THE ELEMENTS OF PERIODIC TABLE
DETAILED DESCRIPTION OF ALL THE ELEMENTS WITH HIGH RESOLUTION PICTURES
BY M. MURUGESH
Neutronium
1
Neutronium 0
N/A ← neutronium → hydrogen
— ↑
Nt [not official] ↓ periodic table - extended periodic table
He
General Name, Symbol, Number
neutronium, Nt [not official], 0
Element category
none
Standard atomic weight
[1] g·mol
Electron configuration
none
−1
Electrons per shell
0 Physical properties
Phase
unknown Miscellaneous Most-stable isotopes
Main article: Isotopes of neutronium iso
NA
1
?
2
?
Nt Nt
half-life 885.7 ± 0.8 s transitory
DM
DE (MeV)
DP
β
−
unknown
1
β−
unknown
2
H H
References
Neutronium is a hypothetical extremely dense phase of matter. The term was originally used in science fiction and in popular literature to refer to a highly dense phase of matter composed primarily of neutrons. The word was coined by scientist Andreas von Antropoff in 1926 (i.e. before the discovery of the neutron itself) for the conjectured 'element of atomic number zero' that he placed at the head of the periodic table.[1] [2] However, the meaning of the term has changed over time, and from the last half of the 20th century onward it has been used legitimately to refer to extremely dense phases of matter resembling the neutron-degenerate matter postulated to exist in the cores of neutron stars.
Neutronium
Neutronium and neutron stars The term neutronium is used in popular literature to refer to the material present in the cores of neutron stars (stars which are too massive to be supported by electron degeneracy pressure and which collapse into a denser phase of matter). This term is very rarely used in scientific literature, for two reasons: • There is no universally agreed-upon definition for the term "neutronium". • There is considerable uncertainty over the composition of the material in the cores of neutron stars (it could be neutron-degenerate matter, strange matter, quark matter, or a variant or combination of the above). When neutron star core material is presumed to consist mostly of free neutrons, it is typically referred to as neutron-degenerate matter in scientific literature.
Neutronium and the periodic table The term neutronium was coined in 1926 by Professor Andreas von Antropoff for a conjectured form of matter made up of neutrons with no protons, which he placed as the chemical element of atomic number zero at the head of his new version of the periodic table. It was subsequently placed as a noble gas in the middle of several spiral representations of the periodic system for classifying the chemical elements, such as the Chemical Galaxy (2005). Although the term is not used in the scientific literature either for a condensed form of matter, or as an element, there have been reports that, besides the free neutron, there may exist two bound forms of neutrons without protons.[3] However, these reports have not been further substantiated. Further information can be found in the following articles: • Mononeutron: Isolated neutrons undergo beta decay with a half-life of approximately 15 minutes, becoming protons (the nucleus of hydrogen), electrons and antineutrinos. • Dineutron: The dineutron, containing two neutrons, is not a bound particle, but has been proposed as an extremely short-lived state produced by nuclear reactions involving tritium. • Trineutron: A trineutron state consisting of three bound neutrons has not been detected, and is not expected to exist even for a short time. • Tetraneutron: A tetraneutron is a hypothetical particle consisting of four bound neutrons. Reports of its existence have not been replicated. If confirmed, it would require revision of current nuclear models.[4] [5] • Pentaneutron: Calculations indicate that the hypothetical pentaneutron state, consisting of a cluster of five neutrons, would not be bound. • And so on, through the numbers, up to icosaneutron, with 20 neutrons.[6] If one accepts neutronium to be an element, the above mentioned neutron clusters would be the isotopes of that element.
Neutronium in fiction The term neutronium has been popular in science fiction since at least the middle of the 20th century. It typically refers to an extremely dense, incredibly strong form of matter. While presumably inspired by the concept of neutron-degenerate matter in the cores of neutron stars, the material used in fiction bears at most only a superficial resemblance, usually depicted as an extremely strong solid under Earth-like conditions, or possessing
2
Neutronium exotic properties such as the ability to manipulate time and space. In contrast, all proposed forms of neutron star core material are fluids and are extremely unstable at pressures lower than that found in stellar cores. Noteworthy appearances of neutronium in fiction include the following: • In Hal Clement's short story Proof (1942), neutronium is the only form of solid matter known to Solarians, the inhabitants of the Sun's interior. • In Vladimir Savchenko's Black Stars (1960), neutronium is mechanically and thermally indestructible substance. It is also used to make antimatter, which leads to a fusion explosion accident. • In Doctor Who (1963), neutronium is a substance which can shield spaces from time-shear when used as shielding in time-vessels. • In Larry Niven's Known Space fictional universe (1964), neutronium is actual neutron star core material. Niven does not make assumptions about its strength, but imagines that small blobs of it would remain stable (and inevitably spherical) under their own gravity. • In the Star Trek universe, neutronium is an extremely hard and durable substance, often used as armor, which conventional weapons cannot penetrate or even dent. • In the computer games Master of Orion (1993), Master of Orion 2 (1996), and Sid Meier's Alpha Centauri (1999), neutronium is the strongest armor type that can be researched in MoO1 and MoO2, and the third strongest in SMAC. MoO1 and MoO2 also feature "neutronium bombs", which are extremely powerful planetary bombardment weapons which causes damage due to gravitic effects. • In Peter F. Hamilton's novel The Neutronium Alchemist (1997), neutronium is created by the "aggressive" setting of a superweapon. • In Stargate SG-1 (1997), neutronium is a substance which is the basis of the technology of the advanced Asgard race, as well as a primary component of human-form Replicators.
See also • Neutron star • Degenerate matter • Neutron-degenerate matter • Compact star
Bibliography • Norman K. Glendenning, R. Kippenhahn, I. Appenzeller, G. Borner, M. Harwit (2000). Compact Stars (2nd ed.).
3
Neutronium
References [1] von Antropoff, A. (1926). " Eine neue Form des periodischen Systems der Elementen. (http:/ / www3. interscience. wiley. com/ cgi-bin/ fulltext/ 112256618/ PDFSTART)" (PDF). Z. Angew. Chem. 39 (23): 722–725. doi: 10.1002/ange.19260392303 (http:/ / dx. doi. org/ 10. 1002/ ange. 19260392303). . Retrieved on 2007-12-12. [2] Stewart, Philip J. (October 2007). " A century on from Dmitrii Mendeleev: tables and spirals, noble gases and Nobel prizes (http:/ / www. springerlink. com/ content/ 6503n26633601877/ )". Foundations of Chemistry 9 (3): 235–245. doi: 10.1007/s10698-007-9038-x (http:/ / dx. doi. org/ 10. 1007/ s10698-007-9038-x). . Retrieved on 2007-12-12. [3] Timofeyuk, N. K. (2003). "Do multineutrons exist?". arΧiv: nucl-th/0301020 (http:/ / www. arxiv. org/ abs/ nucl-th/ 0301020) [nucl-th]. [4] Bertulani, C. A.; Zelevinsky, V. (2002). "Is the tetraneutron a bound dineutron-dineutron molecule?". arΧiv: nucl-th/0212060 (http:/ / www. arxiv. org/ abs/ nucl-th/ 0212060) [nucl-th]. [5] Timofeyuk, N. K. (2002). "On the existence of a bound tetraneutron". arΧiv: nucl-th/0203003 (http:/ / www. arxiv. org/ abs/ nucl-th/ 0203003) [nucl-th]. [6] Bevelacqua, J. J. (June 11, 1981). " Particle stability of the pentaneutron (http:/ / www. sciencedirect. com/ science?_ob=ArticleURL& _udi=B6TVN-472K3HG-2K1& _coverDate=06/ 11/ 1981& _alid=349075295& _rdoc=1& _fmt=& _orig=search& _qd=1& _cdi=5539& _sort=d& view=c& _acct=C000050221& _version=1& _urlVersion=0& _userid=10& md5=f052b79209dd914c85a1bc0d32f774ab)". Physics Letters B 102 (2–3): 79–80. doi: 10.1016/0370-2693(81)91033-9 (http:/ / dx. doi. org/ 10. 1016/ 0370-2693(81)91033-9). .
4
Article Sources and Contributors
Article Sources and Contributors Neutronium Source: http://en.wikipedia.org/w/index.php?oldid=305188587 Contributors: 130.94.122.xxx, 24.93.53.xxx, 3Juno3, Aaryna, Acroterion, Alan Peakall, Anchovee, Arcadian, Arthur Rubin, Ayeroxor, Betacommand, BillC, Bkell, Brighterorange, Bryan Derksen, Cacycle, Caesura, Chairboy, ChrisO, Christopher Thomas, Conversion script, CosineKitty, Cyberia23, Daran, David Latapie, David R. Ingham, Dillee1, Dirac66, Dorftrottel, Dsmith77, Eequor, Evgeny, Gaius Cornelius, Gavinmcq, Geregen2, Gurch, Happy8, Headbomb, Hqb, Hyuri, Iggy Koopa, Ilyak, JHFTC, Jeff G., John Darrow, Joriki, Jwissick, KapilTagore, Keenan Pepper, Kkmurray, Klaxton, Leon..., Ljofa, Looxix, Mac Davis, Manning Bartlett, Mark Foskey, MarkS, Melchoir, MementoVivere, Netizen, Nightscream, Nik42, Nonagonal Spider, OS2Warp, OlEnglish, Omegatron, Oracle7168, Osssua, Pakaran, Pearle, Philip Trueman, PierreAbbat, Prottos007, RJFJR, Rbj, Reyk, Rjwilmsi, Robo37, SFH, Salsb, ScienceApologist, Seminumerical, Shenme, Sobolewski, SocratesJedi, Someguy1221, Soumyasch, Spacepotato, Spartaz, That Guy, From That Show!, The Anome, The Great Attractor, Themel, Tlesinski, Tobyk777, TomTheHand, Trelvis, Urhixidur, Yamamoto Ichiro, 189 anonymous edits
Image Sources, Licenses and Contributors Image:-TableImage.svg Source: http://en.wikipedia.org/w/index.php?title=File:-TableImage.svg License: GNU Free Documentation License Contributors: Bastique, Bibi Saint-Pol, Kwamikagami, Lantrix, Mdd4696, Popolon, Soeb, Tietew, 1 anonymous edits
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5
Hydrogen
1
Hydrogen neutronium ← hydrogen → heliume− ↑ H ↓ Li
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
1H Periodic table
Appearance colorless gas General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensityMelting pointBoiling pointTriple pointCritical pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
at T/K
10 k
100 k
15
20
Atomic properties Oxidation states ElectronegativityIonization energiesCovalent radiusVan der Waals radius Miscellaneous Crystal structureMagnetic orderingThermal conductivitySpeed of soundCAS registry number Most stable isotopes Main article: Isotopes of hydrogen iso
N.A.
half-life
1
99.985%
1
2
0.015%
2
3
trace
12.32 y
H H H
DM
DE (MeV)
DP
H is stable with 0 neutron H is stable with 1 neutron β−
0.01861
3
He
Hydrogen hydrogen, H, 1 nonmetal1, 1, s1.00794(7) g·mol−1 1s1 1 (Image) gas (0 °C, 101.325 kPa) 0.08988 g/L 14.01 K,−259.14 °C,−434.45 °F 20.28 K,−252.87 °C,−423.17 °F 13.8033 K (-259°C), 7.042 kPa 32.97 K, 1.293 MPa (H2) 0.117 kJ·mol−1 (H2) 0.904 kJ·mol−1 (25 °C) (H2) −1 −1 28.836 J·mol ·K 1, −1 (amphoteric oxide) 2.20 (Pauling scale) 1st: 1312.0 kJ·mol−131±5 pm 120 pm hexagonal diamagnetic[1] (300 K) 180.5 m W·m−1·K−1 (gas, 27 °C) 1310 m/s 1333-74-0 Hydrogen (pronounced /ˈhaɪdrədʒən/[2] ) is the chemical element with atomic number 1. It is represented by the symbol H. At standard temperature and pressure, hydrogen is a colorless, odorless, nonmetallic, tasteless, highly flammable diatomic gas with the molecular formula H2. With an atomic weight of , hydrogen is the lightest element. Hydrogen is the most abundant chemical element, constituting roughly 75% of the universe's elemental mass.[3] Stars in the main sequence are mainly composed of hydrogen in its plasma state. Elemental hydrogen is relatively rare on Earth. Industrial production is from hydrocarbons such as methane with most being used "captively" at the production site. The two largest uses are in fossil fuel processing (e.g., hydrocracking) and ammonia production mostly for the fertilizer market. Hydrogen may be produced from the electrolysis of water or other hydrogen production methods like the reforming of natural gas.[4] The most common isotope of hydrogen is protium (name rarely used, symbol H) with a single proton and no neutrons. In ionic compounds it can take a negative charge (an anion known as a hydride and written as H−), or as a positively-charged species H+. The latter cation is written as though composed of a bare proton, but in reality, hydrogen cations in ionic compounds always occur as more complex species. Hydrogen forms compounds with most elements and is present in water and most organic compounds. It plays a particularly important role in acid-base chemistry with many reactions exchanging protons between soluble molecules. As the only neutral atom with an analytic solution to the Schrödinger equation, the study of the energetics and bonding of the hydrogen atom played a key role in the development of quantum mechanics. Hydrogen is important in metallurgy as it can embrittle many metals,[5] complicating the design of pipelines and storage tanks.[6] Hydrogen is highly soluble in many rare earth and transition metals[7] and is soluble in both nanocrystalline and amorphous metals.[8] Hydrogen solubility in metals is influenced by local distortions or impurities in the crystal lattice.[9]
2
Hydrogen
3
Combustion Hydrogen gas (dihydrogen[10] ) is highly flammable and will burn in air at a very wide range of concentrations between 4% and 75% by volume.[11] The enthalpy of combustion for hydrogen is −286 kJ/mol:[12] 2 H2(g) + O2(g) → 2 H2O(l) + 572 kJ (286 kJ/mol)[13] Hydrogen-oxygen mixtures are explosive across a wide range of proportions. Its autoignition temperature, the temperature at which it ignites spontaneously in air, is 560 °C (1040 °F).[14] Pure hydrogen-oxygen flames emit ultraviolet light and are The Space Shuttle Main nearly invisible to the naked eye as illustrated by the faint Engine burns hydrogen with plume of the Space Shuttle main engine compared to the highly oxygen, producing a nearly visible plume of a Space Shuttle Solid Rocket Booster). The invisible flame detection of a burning hydrogen leak may require a flame detector; such leaks can be very dangerous. The explosion of the Hindenburg airship was an infamous example of hydrogen combustion; the cause is debated, but the visible flames were the result of combustible materials in the ship's skin.[15] Because hydrogen is buoyant in air, hydrogen flames tend to ascend rapidly and cause less damage than hydrocarbon fires. Two-thirds of the Hindenburg passengers survived the fire, and many deaths were instead the result of falls or burning diesel fuel.[16] H2 reacts with every oxidizing element. Hydrogen can react spontaneously and violently at room temperature with chlorine and fluorine to form the corresponding halides: hydrogen chloride and hydrogen fluoride.[17]
Electron energy levels The ground state energy level of the electron in a hydrogen atom is −13.6 eV, which is equivalent to an ultraviolet photon of roughly 92 nm.[18] The energy levels of hydrogen can be calculated fairly accurately using the Bohr model of the atom, which conceptualizes the electron as "orbiting" the proton in analogy to the Earth's orbit of the sun. However, the electromagnetic force attracts electrons and protons to one another, while planets and celestial objects are attracted to each other by gravity. Because of the discretization of angular momentum postulated in early quantum mechanics by Bohr, the electron in the Bohr model can only occupy certain allowed distances from the proton, and therefore only certain allowed energies.[19]
Depiction of a hydrogen atom showing the diameter as about twice the Bohr model radius (image not to scale).
A more accurate description of the hydrogen atom comes from a purely quantum mechanical treatment that uses the Schrödinger equation or the equivalent Feynman path integral formulation to calculate the probability density of the electron around the proton.[20]
Hydrogen
Elemental molecular forms There exist two different spin isomers of hydrogen diatomic molecules that differ by the relative spin of their nuclei.[21] In the orthohydrogen form, the spins of the two protons are parallel and form a triplet state; in the parahydrogen form the spins are antiparallel and form a singlet. At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form, also known as the "normal form".[22] The equilibrium ratio of orthohydrogen to parahydrogen depends on temperature, but since the ortho form is an excited state and has a higher energy than the para form, it is unstable and cannot be purified. At very low temperatures, the equilibrium state is composed almost exclusively of the para form. The liquid and gas phase thermal properties of pure First tracks observed in liquid parahydrogen differ significantly from those of the hydrogen bubble chamber at the normal form because of differences in rotational heat Bevatron capacities, as discussed more fully in Spin isomers of hydrogen.[23] The ortho/para distinction also occurs in other hydrogen-containing molecules or functional groups, such as water and methylene, but is of little significance with respect to thermal properties.[24] The uncatalyzed interconversion between para and ortho H2 increases with increasing temperature; thus rapidly condensed H2 contains large quantities of the high-energy ortho form that convert to the para form very slowly.[25] The ortho/para ratio in condensed H2 is an important consideration in the preparation and storage of liquid hydrogen: the conversion from ortho to para is exothermic and produces enough heat to evaporate the hydrogen liquid, leading to loss of the liquefied material. Catalysts for the ortho-para interconversion, such as ferric oxide, activated carbon, platinized asbestos, rare earth metals, uranium compounds, chromic oxide, or some nickel[26] compounds, are used during hydrogen cooling.[27] A molecular form called protonated molecular hydrogen, or H+3, is found in the interstellar medium (ISM), where it is generated by ionization of molecular hydrogen from cosmic rays. It has also been observed in the upper atmosphere of the planet Jupiter. This molecule is relatively stable in the environment of outer space due to the low temperature and density. H+3 is one of the most abundant ions in the Universe, and it plays a notable role in the chemistry of the interstellar medium.[28]
Compounds Covalent and organic compounds While H2 is not very reactive under standard conditions, it does form compounds with most elements. Millions of hydrocarbons are known, but they are not formed by the direct reaction of elementary hydrogen and carbon. Hydrogen can form compounds with elements that are more electronegative, such as halogens (e.g., F, Cl, Br, I); in these compounds
4
Hydrogen hydrogen takes on a partial positive charge.[29] When bonded to fluorine, oxygen, or nitrogen, hydrogen can participate in a form of strong noncovalent bonding called hydrogen bonding, which is critical to the stability of many biological molecules.[30] [31] Hydrogen also forms compounds with less electronegative elements, such as the metals and metalloids, in which it takes on a partial negative charge. These compounds are often known as hydrides.[32] Hydrogen forms a vast array of compounds with carbon. Because of their general association with living things, these compounds came to be called organic compounds;[33] the study of their properties is known as organic chemistry[34] and their study in the context of living organisms is known as biochemistry.[35] By some definitions, "organic" compounds are only required to contain carbon. However, most of them also contain hydrogen, and since it is the carbon-hydrogen bond which gives this class of compounds most of its particular chemical characteristics, carbon-hydrogen bonds are required in some definitions of the word "organic" in chemistry.[33] In inorganic chemistry, hydrides can also serve as bridging ligands that link two metal centers in a coordination complex. This function is particularly common in group 13 elements, especially in boranes (boron hydrides) and aluminium complexes, as well as in clustered carboranes.[36]
Hydrides Compounds of hydrogen are often called hydrides, a term that is used fairly loosely. The term "hydride" implies that the H atom has acquired a negative or anionic character, denoted H−, and is used when hydrogen forms a compound with a more electropositive element. The existence of the hydride anion, suggested by Gilbert N. Lewis in 1916 for group I and II salt-like hydrides, was demonstrated by Moers in 1920 with the electrolysis of molten lithium hydride (LiH), that produced a stoichiometric quantity of hydrogen at the anode.[37] For hydrides other than group I and II metals, the term is quite misleading, considering the low electronegativity of hydrogen. An exception in group II hydrides is BeH2, which is polymeric. In lithium aluminium hydride, the AlH−4 anion carries hydridic centers firmly attached to the Al(III). Although hydrides can be formed with almost all main-group elements, the number and combination of possible compounds varies widely; for example, there are over 100 binary borane hydrides known, but only one binary aluminium hydride.[38] Binary indium hydride has not yet been identified, although larger complexes exist.[39]
Protons and acids Oxidation of hydrogen, in the sense of removing its electron, formally gives H+, containing no electrons and a nucleus which is usually composed of one proton. That is why H+ is often called a proton. This species is central to discussion of acids. Under the Bronsted-Lowry theory, acids are proton donors, while bases are proton acceptors. A bare proton, H+, cannot exist in solution or in ionic crystals, because of its unstoppable attraction to other atoms or molecules with electrons. Except at the high temperatures associated with plasmas, such protons cannot be removed from the electron clouds of atoms and molecules, and will remain attached to them. However, the term 'proton' is sometimes used loosely and metaphorically to refer to positively charged or cationic hydrogen attached to other species in this fashion, and as such is denoted "H+" without any
5
Hydrogen implication that any single protons exist freely as a species. To avoid the implication of the naked "solvated proton" in solution, acidic aqueous solutions are sometimes considered to contain a less unlikely fictitious species, termed the "hydronium ion" (H3O+). However, even in this case, such solvated hydrogen cations are thought more realistically physically to be organized into clusters that form species closer to H9O+4.[40] Other oxonium ions are found when water is in solution with other solvents.[41] Although exotic on earth, one of the most common ions in the universe is the H3+ ion, known as protonated molecular hydrogen or the triatomic hydrogen cation.[42]
Isotopes Hydrogen has three naturally occurring isotopes, denoted 1H, 2H and 3H. Other, highly unstable nuclei (4H to 7H) have been synthesized in the laboratory but not observed in nature.[43] [44] •
1
H is the most common hydrogen isotope with an
abundance of more than 99.98%. Because the nucleus of this isotope consists of only a single proton, it is given the descriptive but rarely used formal name protium.[45] •
2
H, the other stable hydrogen isotope, is known as
deuterium and contains one proton and one neutron Protium, the most common isotope of in its nucleus. Essentially all deuterium in the hydrogen, has one proton and one universe is thought to have been produced at the electron. Unique among all stable time of the Big Bang, and has endured since that isotopes, it has no neutrons (see time. Deuterium is not radioactive, and does not diproton for discussion of why others do not exist). represent a significant toxicity hazard. Water enriched in molecules that include deuterium instead of normal hydrogen is called heavy water. Deuterium and its compounds are used as a non-radioactive label in chemical experiments and in solvents for 1H-NMR spectroscopy.[46] Heavy water is used as a neutron moderator and coolant for nuclear reactors. Deuterium is also a potential fuel for commercial nuclear fusion.[47] •
3
H is known as tritium and contains one proton and two neutrons in its nucleus. It is radioactive, decaying into Helium-3 through beta decay with a half-life of 12.32 years.[36] Small amounts of tritium occur naturally because of the interaction of cosmic rays with atmospheric gases; tritium has also been released during nuclear weapons tests.[48] It is used in nuclear fusion reactions,[49] as a tracer in isotope geochemistry,[50] and specialized in self-powered lighting devices.[51] Tritium has also been used in chemical and biological labeling experiments as a radiolabel.[52]
Hydrogen is the only element that has different names for its isotopes in common use today. (During the early study of radioactivity, various heavy radioactive isotopes were given names, but such names are no longer used). The symbols D and T (instead of 2H and 3 H) are sometimes used for deuterium and tritium, but the corresponding symbol P is already in use for phosphorus and thus is not available for protium.[53] In its nomenclatural guidelines, the International Union of Pure and Applied Chemistry allows any of D, T, 2H,
6
Hydrogen
7
and 3H to be used, although 2H and 3H are preferred.[54]
Natural occurrence Hydrogen is the most abundant element in the universe, making up 75% of normal matter by mass and over 90% by number of atoms.[55] This element is found in great abundance in stars and gas giant planets. Molecular clouds of H2 are associated with star formation. Hydrogen plays a vital role in powering stars through proton-proton reaction and CNO cycle nuclear fusion.[56] Throughout the universe, hydrogen is mostly found in the atomic and plasma states whose properties are quite different from molecular hydrogen. As a plasma, hydrogen's electron and proton are not bound together, NGC 604, a giant region of ionized resulting in very high electrical conductivity and high hydrogen in the Triangulum Galaxy emissivity (producing the light from the sun and other stars). The charged particles are highly influenced by magnetic and electric fields. For example, in the solar wind they interact with the Earth's magnetosphere giving rise to Birkeland currents and the aurora. Hydrogen is found in the neutral atomic state in the Interstellar medium. The large amount of neutral hydrogen found in the damped Lyman-alpha systems is thought to dominate the cosmological baryonic density of the Universe up to redshift z=4.[57] Under ordinary conditions on Earth, elemental hydrogen exists as the diatomic gas, H2 (for data see table). However, hydrogen gas is very rare in the Earth's atmosphere (1 ppm by volume) because of its light weight, which enables it to escape from Earth's gravity more easily than heavier gases. However, hydrogen (in chemically combined form) is the third most abundant element on the Earth's surface.[58] Most of the Earth's hydrogen is in the form of chemical compounds such as hydrocarbons and water.[36] Hydrogen gas is produced by some bacteria and algae and is a natural component of flatus. Methane is a hydrogen source of increasing importance.[59]
History Discovery and use Hydrogen gas, H2, was first artificially produced and formally described by T. Von Hohenheim (also known as Paracelsus, 1493–1541) via the mixing of metals with strong acids.[60] He was unaware that the flammable gas produced by this chemical reaction was a new chemical element. In 1671, Robert Boyle rediscovered and described the reaction between iron filings and dilute acids, which results in the production of hydrogen gas.[61] In 1766, Henry Cavendish was the first to recognize hydrogen gas as a discrete substance, by identifying the gas from a metal-acid reaction as "inflammable air" and further finding in 1781 that the gas produces water when burned. He is usually given credit for its discovery as an element.[62] [63] In 1783, Antoine Lavoisier gave the element the name hydrogen (from the Greek hydro meaning water and genes meaning creator)[64] when he and Laplace
Hydrogen reproduced Cavendish's finding that water is produced when hydrogen is burned.[63] Hydrogen was liquefied for the first time by James Dewar in 1898 by using regenerative cooling and his invention, the vacuum flask.[63] He produced solid hydrogen the next year.[63] Deuterium was discovered in December 1931 by Harold Urey, and tritium was prepared in 1934 by Ernest Rutherford, Mark Oliphant, and Paul Harteck.[62] Heavy water, which consists of deuterium in the place of regular hydrogen, was discovered by Urey's group in 1932.[63] François Isaac de Rivaz built the first internal combustion engine powered by a mixture of hydrogen and oxygen in 1806. Edward Daniel Clarke invented the hydrogen gas blowpipe in 1819. The Döbereiner's lamp and limelight were invented in 1823.[63] The first hydrogen-filled balloon was invented by Jacques Charles in 1783.[63] Hydrogen provided the lift for the first reliable form of air-travel following the 1852 invention of the first hydrogen-lifted airship by Henri Giffard.[63] German count Ferdinand von Zeppelin promoted the idea of rigid airships lifted by hydrogen that later were called Zeppelins; the first of which had its maiden flight in 1900.[63] Regularly scheduled flights started in 1910 and by the outbreak of World War I in August 1914, they had carried 35,000 passengers without a serious incident. Hydrogen-lifted airships were used as observation platforms and bombers during the war. The first non-stop transatlantic crossing was made by the British airship R34 in 1919. Regular passenger service resumed in the 1920s and the discovery of helium reserves in the United States promised increased safety, but the U.S. government refused to sell the gas for this purpose. Therefore, H2 was used in the Hindenburg airship, which was destroyed in a midair fire over New Jersey on May 6, 1937.[63] The incident was broadcast live on radio and filmed. Ignition of leaking hydrogen as widely assumed to be the cause but later investigations pointed to ignition of the aluminized fabric coating by static electricity. But the damage to hydrogen's reputation as a lifting gas was already done. In the same year the first hydrogen-cooled turbogenerator went into service with gaseous hydrogen as a coolant in the rotor and the stator in 1937 at Dayton, Ohio, by the Dayton Power & Light Co,[65] because of the thermal conductivity of hydrogen gas this is the most common type in its field today. The nickel hydrogen battery was used for the first time in 1977 aboard the U.S. Navy's Navigation technology satellite-2 (NTS-2).[66] For example, the ISS,[67] Mars Odyssey[68] and the Mars Global Surveyor[69] are equipped with nickel-hydrogen batteries. The Hubble Space Telescope, at the time its original batteries were finally changed in May 2009, more than 19 years after launch, led with the highest number of charge/discharge cycles of any NiH2 battery in low earth orbit.[70]
Role in quantum theory Because of its relatively simple atomic structure, consisting only of a proton and an electron, the hydrogen atom, Hydrogen emission spectrum lines in the visible range. These are the four visible lines of the Balmer series together with the spectrum of light produced from it or absorbed by it, has been central to the development of the theory of atomic structure.[71] Furthermore, the corresponding simplicity of the hydrogen molecule and the corresponding cation H2+ allowed fuller understanding of the nature of the chemical bond, which followed
8
Hydrogen shortly after the quantum mechanical treatment of the hydrogen atom had been developed in the mid-1920s. One of the first quantum effects to be explicitly noticed (but not understood at the time) was a Maxwell observation involving hydrogen, half a century before full quantum mechanical theory arrived. Maxwell observed that the specific heat capacity of H2 unaccountably departs from that of a diatomic gas below room temperature and begins to increasingly resemble that of a monatomic gas at cryogenic temperatures. According to quantum theory, this behavior arises from the spacing of the (quantized) rotational energy levels, which are particularly wide-spaced in H2 because of its low mass. These widely spaced levels inhibit equal partition of heat energy into rotational motion in hydrogen at low temperatures. Diatomic gases composed of heavier atoms do not have such widely spaced levels and do not exhibit the same effect.[72]
Production H2 is produced in chemistry and biology laboratories, often as a by-product of other reactions; in industry for the hydrogenation of unsaturated substrates; and in nature as a means of expelling reducing equivalents in biochemical reactions.
Laboratory In the laboratory, H2 is usually prepared by the reaction of acids on metals such as zinc with Kipp's apparatus. Zn + 2 H+ → Zn2+ + H2 Aluminium can also produce H2 upon treatment with bases: 2 Al + 6 H2O + 2 OH− → 2 Al(OH)4− + 3 H2 The electrolysis of water is a simple method of producing hydrogen. A low voltage current is run through the water, and gaseous oxygen forms at the anode while gaseous hydrogen forms at the cathode. Typically the cathode is made from platinum or another inert metal when producing hydrogen for storage. If, however, the gas is to be burnt on site, oxygen is desirable to assist the combustion, and so both electrodes would be made from inert metals. (Iron, for instance, would oxidize, and thus decrease the amount of oxygen given off.) The theoretical maximum efficiency (electricity used vs. energetic value of hydrogen produced) is between 80–94%.[73] 2H2O(aq) → 2H2(g) + O2(g) In 2007, it was discovered that an alloy of aluminium and gallium in pellet form added to water could be used to generate hydrogen. The process also creates alumina, but the expensive gallium, which prevents the formation of an oxide skin on the pellets, can be re-used. This has important potential implications for a hydrogen economy, since hydrogen can be produced on-site and does not need to be transported.[74]
Industrial Hydrogen can be prepared in several different ways, but economically the most important processes involve removal of hydrogen from hydrocarbons. Commercial bulk hydrogen is usually produced by the steam reforming of natural gas.[75] At high temperatures (1000–1400 K, °C;700–1100 °C or 1,300–2,000 °F), steam (water vapor) reacts with methane to yield carbon monoxide and H2.
9
Hydrogen CH4 + H2O → CO + 3 H2 This reaction is favored at low pressures but is nonetheless conducted at high pressures (2.0 MPa, 20 atm or 600 inHg) since high pressure H2 is the most marketable product. The product mixture is known as "synthesis gas" because it is often used directly for the production of methanol and related compounds. Hydrocarbons other than methane can be used to produce synthesis gas with varying product ratios. One of the many complications to this highly optimized technology is the formation of coke or carbon: CH4 → C + 2 H2 Consequently, steam reforming typically employs an excess of H2O. Additional hydrogen can be recovered from the steam by use of carbon monoxide through the water gas shift reaction, especially with an iron oxide catalyst. This reaction is also a common industrial source of carbon dioxide:[75] CO + H2O → CO2 + H2 Other important methods for H2 production include partial oxidation of hydrocarbons:[76] 2 CH4 + O2 → 2 CO + 4 H2 and the coal reaction, which can serve as a prelude to the shift reaction above:[75] C + H2O → CO + H2 Hydrogen is sometimes produced and consumed in the same industrial process, without being separated. In the Haber process for the production of ammonia, hydrogen is generated from natural gas.[77] Electrolysis of brine to yield chlorine also produces hydrogen as a co-product.[78]
Thermochemical There are more than 200 thermochemical cycles which can be used for water splitting, around a dozen of these cycles such as the iron oxide cycle, cerium(IV) oxide-cerium(III) oxide cycle, zinc zinc-oxide cycle, sulfur-iodine cycle, copper-chlorine cycle and hybrid sulfur cycle are under research and in testing phase to produce hydrogen and oxygen from water and heat without using electricity.[79] A number of laboratories (including in France, Germany, Greece, Japan, and the USA) are developing thermochemical methods to produce hydrogen from solar energy and water.[80]
Applications Large quantities of H2 are needed in the petroleum and chemical industries. The largest application of H2 is for the processing ("upgrading") of fossil fuels, and in the production of ammonia. The key consumers of H2 in the petrochemical plant include hydrodealkylation, hydrodesulfurization, and hydrocracking. H2 has several other important uses. H2 is used as a hydrogenating agent, particularly in increasing the level of saturation of unsaturated fats and oils (found in items such as margarine), and in the production of methanol. It is similarly the source of hydrogen in the manufacture of hydrochloric acid. H2 is also used as a reducing agent of metallic ores.[81] Apart from its use as a reactant, H2 has wide applications in physics and engineering. It is used as a shielding gas in welding methods such as atomic hydrogen welding.[82] [83] H2 is used as the rotor coolant in electrical generators at power stations, because it has the highest thermal conductivity of any gas. Liquid H2 is used in cryogenic research, including
10
Hydrogen
11
superconductivity studies.[84] Since H2 is lighter than air, having a little more than [85]
1
⁄15 of
the density of air, it was once widely used as a lifting gas in balloons and airships.
In more recent applications, hydrogen is used pure or mixed with nitrogen (sometimes called forming gas) as a tracer gas for minute leak detection. Applications can be found in the automotive, chemical, power generation, aerospace, and telecommunications industries.[86] Hydrogen is an authorized food additive (E 949) that allows food package leak testing among other anti-oxidizing properties.[87] Hydrogen's rarer isotopes also each have specific applications. Deuterium (hydrogen-2) is used in nuclear fission applications as a moderator to slow neutrons, and in nuclear fusion reactions.[63] Deuterium compounds have applications in chemistry and biology in studies of reaction isotope effects.[88] Tritium (hydrogen-3), produced in nuclear reactors, is used in the production of hydrogen bombs,[89] as an isotopic label in the biosciences,[52] and as a radiation source in luminous paints.[90] The triple point temperature of equilibrium hydrogen is a defining fixed point on the ITS-90 temperature scale at 13.8033 kelvins.[91]
Energy carrier Hydrogen is not an energy resource,[92] except in the hypothetical context of commercial nuclear fusion power plants using deuterium or tritium, a technology presently far from development.[93] The Sun's energy comes from nuclear fusion of hydrogen, but this process is difficult to achieve controllably on Earth.[94] Elemental hydrogen from solar, biological, or electrical sources require more energy to make it than is obtained by burning it, so in these cases hydrogen functions as an energy carrier, like a battery. Hydrogen may be obtained from fossil sources (such as methane), but these sources are unsustainable.[92] The energy density per unit volume of both liquid hydrogen and compressed hydrogen gas at any practicable pressure is significantly less than that of traditional fuel sources, although the energy density per unit fuel mass is higher.[92] Nevertheless, elemental hydrogen has been widely discussed in the context of energy, as a possible future carrier of energy on an economy-wide scale.[95] For example, CO2 sequestration followed by carbon capture and storage could be conducted at the point of H2 production from fossil fuels.[96] Hydrogen used in transportation would burn relatively cleanly, with some NOx emissions,[97] but without carbon emissions.[96] However, the infrastructure costs associated with full conversion to a hydrogen economy would be substantial.[98]
Biological reactions H2 is a product of some types of anaerobic metabolism and is produced by several microorganisms, usually via reactions catalyzed by iron- or nickel-containing enzymes called hydrogenases. These enzymes catalyze the reversible redox reaction between H2 and its component two protons and two electrons. Creation of hydrogen gas occurs in the transfer of reducing equivalents produced during pyruvate fermentation to water.[99] Water splitting, in which water is decomposed into its component protons, electrons, and oxygen, occurs in the light reactions in all photosynthetic organisms. Some such organisms—including the alga Chlamydomonas reinhardtii and cyanobacteria—have evolved a second step in the dark reactions in which protons and electrons are reduced to form H2 gas by specialized hydrogenases in the chloroplast.[100] Efforts have been
Hydrogen undertaken to genetically modify cyanobacterial hydrogenases to efficiently synthesize H2 gas even in the presence of oxygen.[101] Efforts have also been undertaken with genetically modified alga in a bioreactor.[102]
Safety and precautions Hydrogen poses a number of hazards to human safety, from potential detonations and fires when mixed with air to being an asphyxant in its pure, oxygen-free form.[104] In addition, liquid hydrogen is a cryogen and presents dangers (such as frostbite) associated with very cold liquids.[105] Hydrogen dissolves in some metals, and, in addition to leaking out, may have adverse effects on them, such as hydrogen embrittlement.[106] Hydrogen gas leaking into external air may spontaneously ignite. Moreover, hydrogen fire, while being extremely hot, is almost invisible, and thus can lead to accidental burns.[107] Even interpreting the hydrogen data (including safety data) is confounded by a number of phenomena. Many Inhalation of air with high physical and chemical properties of hydrogen depend concentration of hydrogen as with on the parahydrogen/orthohydrogen ratio (it often takes hydrox and hydreliox displaces oxygen and may cause the above symptoms as days or weeks at a given temperature to reach the [103] an asphyxant. equilibrium ratio, for which the data is usually given). Hydrogen detonation parameters, such as critical detonation pressure and temperature, strongly depend on the container geometry.[104]
See also • Antihydrogen • • • • • • • • • • • • •
Blacklight Power Hydrogen cycle Hydrogen leak testing Hydrogen-like atom Hydrogen line Hydrogen planes Hydrogen spectral series Hydrogen station Hydrogen technologies Hydrogen vehicle Metallic hydrogen Oxyhydrogen Photohydrogen
12
Hydrogen
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" New process generates hydrogen from aluminum alloy to run engines, fuel cells (http:/ / news. uns. purdue. edu/ x/ 2007a/ 070515WoodallHydrogen. html)". Purdue University. . Retrieved 2008-02-05. [75] Oxtoby, D. W. (2002). Principles of Modern Chemistry (5th edition ed.). Thomson Brooks/Cole. ISBN 0030353734. [76] " Hydrogen Properties, Uses, Applications (http:/ / www. uigi. com/ hydrogen. html)". Universal Industrial Gases, Inc.. 2007. . Retrieved 2008-03-11. [77] Funderburg, Eddie (2008). " Why Are Nitrogen Prices So High? (http:/ / www. noble. org/ Ag/ Soils/ NitrogenPrices/ Index. htm)". The Samuel Roberts Noble Foundation. . Retrieved 2008-03-11. [78] Lees, Andrew (2007). " Chemicals from salt (http:/ / www. bbc. co. uk/ schools/ gcsebitesize/ chemistry/ usefulproductsrocks/ chemicals_saltrev3. shtml)". BBC. . Retrieved 2008-03-11. [79] " Development of solar-powered thermochemical production of hydrogen from water (http:/ / www. hydrogen. energy. gov/ pdfs/ review05/ pd28_weimer. pdf)" (PDF). . [80] Robert Perret. " Development of Solar-Powered Thermochemical Production of Hydrogen from Water, DOE Hydrogen Program, 2007 (http:/ / www. hydrogen. energy. gov/ pdfs/ progress07/ ii_f_1_perret. pdf)" (PDF). . Retrieved 2008-05-17. [81] Chemistry Operations (2003-12-15). " Hydrogen (http:/ / periodic. lanl. gov/ elements/ 1. html)". Los Alamos National Laboratory. . Retrieved 2008-02-05. [82] Durgutlu, Ahmet (2003-10-27). " Experimental investigation of the effect of hydrogen in argon as a shielding gas on TIG welding of austenitic stainless steel (http:/ / www. sciencedirect. com/ science?_ob=ArticleURL& _udi=B6TX5-49W1W1V-7& _user=10& _rdoc=1& _fmt=& _orig=search& _sort=d& view=c& _acct=C000050221& _version=1& _urlVersion=0& _userid=10& md5=2074bcd5650e0ca62aa09b8713708226)". ScienceDirect (Ankara, Turkey: Gazi University) 25 (1): 19–23. doi: 10.1016/j.matdes.2003.07.004 (http:/ / dx. doi. org/ 10. 1016/ j. matdes. 2003. 07. 004). . Retrieved 2008-04-06. [83] " Atomic Hydrogen Welding (http:/ / www. specialwelds. com/ underwater-welding/ atomic-hydrogen-welding. htm)". Specialty Welds. 2007. . [84] Hardy, Walter N. (2003-03-19). " From H2 to cryogenic H masers to HiTc superconductors: An unlikely but rewarding path (http:/ / www. sciencedirect. com/ science?_ob=ArticleURL& _udi=B6TVJ-485PG6D-D& _user=10& _rdoc=1& _fmt=& _orig=search& _sort=d& view=c& _acct=C000050221& _version=1& _urlVersion=0& _userid=10& md5=f4ec8a7def03583c043dd9e60aa0c07e)". Physica C: Superconductivity (Vancouver, Canada: University of British Columbia) 388–389: 1–6. doi: 10.1016/S0921-4534(02)02591-1 (http:/ / dx. doi. org/ 10. 1016/ S0921-4534(02)02591-1). . Retrieved 2008-03-25. [85] Barnes, Matthew (2004). " LZ-129, Hindenburg (http:/ / www. ciderpresspottery. com/ ZLA/ greatzeps/ german/ Hindenburg. html)". The Great Zeppelins. . Retrieved 2008-03-18. [86] Block, Matthias (3 September 2004). " Hydrogen as Tracer Gas for Leak Detection (http:/ / www. ndt. net/ abstract/ wcndt2004/ 523. htm)". 16th WCNDT 2004. Montreal, Canada: Sensistor Technologies. [87] " Report from the Commission on Dietary Food Additive Intake (http:/ / ec. europa. eu/ food/ fs/ sfp/ addit_flavor/ flav15_en. pdf)" (PDF). European Union. . Retrieved 2008-02-05. [88] Reinsch, J; A Katz, J Wean, G Aprahamian, JT MacFarland (10 October 1980). " The deuterium isotope effect upon the reaction of fatty acyl-CoA dehydrogenase and butyryl-CoA (http:/ / www. jbc. org/ cgi/ content/ abstract/ 255/ 19/ 9093)". J. Biol. Chem. 255 (19): 9093–97. PMID 7410413. . Retrieved 2008-03-24. [89] Bergeron, Kenneth D. (January–February 2004). " The Death of no-dual-use (http:/ / find. galegroup. com/ itx/ start. do?prodId=SPJ. SP06)". Bulletin of the Atomic Scientists (Educational Foundation for Nuclear Science, Inc.) 60 (1): 15. doi: 10.2968/060001004 (http:/ / dx. doi. org/ 10. 2968/ 060001004). . Retrieved 2008-04-13. [90] Quigg, Catherine T. (March 1984). " Tritium Warning (http:/ / search. ebscohost. com/ login. aspx?direct=true& db=sch& AN=11181317& site=ehost-live)". Bulletin of the Atomic Scientists (Chicago) 40 (3): 56–57. ISSN 0096-3402 (http:/ / worldcat. org/ issn/ 0096-3402). . Retrieved 2008-04-15.
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Hydrogen [91] " International Temperature Scale of 1990 (http:/ / www. bipm. org/ utils/ common/ pdf/ its-90/ ITS-90. pdf)" (PDF). Procès-Verbaux du Comité International des Poids et Mesures. 1989. pp. T23–T42. [92] McCarthy, John (1995-12-31). " Hydrogen (http:/ / www-formal. stanford. edu/ jmc/ progress/ hydrogen. html)". Stanford University. . Retrieved 2008-03-14. [93] " Nuclear Fusion Power (http:/ / www. world-nuclear. org/ info/ inf66. html)". World Nuclear Association. May 2007. . Retrieved 2008-03-16. [94] " Chapter 13: Nuclear Energy — Fission and Fusion (http:/ / www. energyquest. ca. gov/ story/ chapter13. html)". Energy Story. California Energy Commission. 2006. . Retrieved 2008-03-14. [95] US Department of Energy (2006-03-22). " DOE Seeks Applicants for Solicitation on the Employment Effects of a Transition to a Hydrogen Economy (http:/ / www. hydrogen. energy. gov/ news_transition. html)". Press release. . Retrieved 2008-03-16. [96] Georgia Tech (2008-02-11). " Carbon Capture Strategy Could Lead to Emission-Free Cars (http:/ / www. gatech. edu/ newsroom/ release. html?id=1707)". Press release. . Retrieved 2008-03-16. [97] Heffel, James W. (24 December 2002). " NOx emission and performance data for a hydrogen fueled internal combustion engine at 1500 rpm using exhaust gas recirculation (http:/ / www. sciencedirect. com/ science?_ob=ArticleURL& _udi=B6V3F-47HJVY6-8& _user=10& _rdoc=1& _fmt=& _orig=search& _sort=d& view=c& _acct=C000050221& _version=1& _urlVersion=0& _userid=10& md5=bbc8c5bce46f1d4ba3a814f5c828ee80)". International Journal of Hydrogen Energy (Riverside, CA: University of California) 28 (8): 901–908. doi: 10.1016/S0360-3199(02)00157-X (http:/ / dx. doi. org/ 10. 1016/ S0360-3199(02)00157-X). . Retrieved 2008-03-16. [98] See Romm, Joseph J. (2004). The Hype About Hydrogen: Fact And Fiction In The Race To Save The Climate (1st edition ed.). Island Press. ISBN 155963703X. [99] Cammack, Richard; Robson, R. L. (2001). Hydrogen as a Fuel: Learning from Nature. Taylor & Francis Ltd. ISBN 0415242428. [100] Kruse, O.; Rupprecht, J.; Bader, K.-P.; Thomas-Hall, S.; Schenk, P. M.; Finazzi, G.; Hankamer, B (2005). "Improved photobiological H2 production in engineered green algal cells". The Journal of Biological Chemistry 280 (40): 34170–7. doi: 10.1074/jbc.M503840200 (http:/ / dx. doi. org/ 10. 1074/ jbc. M503840200). PMID 16100118. [101] Smith, H. O.; Xu, Q (2005). " IV.E.6 Hydrogen from Water in a Novel Recombinant Oxygen-Tolerant Cyanobacteria System (http:/ / ec. europa. eu/ food/ fs/ sfp/ addit_flavor/ flav15_en. pdf)" (PDF). FY2005 Progress Report. United States Department of Energy. . Retrieved 2008-02-05. [102] Williams, Chris (2006-02-24). " Pond life: the future of energy (http:/ / www. theregister. co. uk/ 2006/ 02/ 24/ pond_scum_breakthrough/ )". Science (The Register). . Retrieved 2008-03-24. [103] Lenntech: Health effects of hydrogen - Environmental effects of hydrogen (http:/ / www. lenntech. com/ Periodic-chart-elements/ H-en. htm) Received on 11 February, 2009 [104] Smith, H. O.; Xu, Q (1997). " Safety Standard for Hydrogen and Hydrogen Systems (http:/ / www. hq. nasa. gov/ office/ codeq/ doctree/ canceled/ 871916. pdf)" (PDF). NASA. . Retrieved 2008-02-05. [105] " Liquid Hydrogen MSDS (http:/ / www. hydrogenandfuelcellsafety. info/ resources/ mdss/ Praxair-LH2. pdf)" (PDF). Praxair, Inc.. September 2004. . Retrieved 2008-04-16. [106] " 'Bugs' and hydrogen embrittlement (http:/ / search. ebscohost. com/ login. aspx?direct=true& db=sch& AN=8837940& site=ehost-live)". Science News (Washington D.C.) 128 (3): 41. 1985-07-20. doi: 10.2307/3970088 (http:/ / dx. doi. org/ 10. 2307/ 3970088). ISSN 0036-8423 (http:/ / worldcat. org/ issn/ 0036-8423). . Retrieved 2008-04-16. [107] " Hydrogen Safety (http:/ / www. humboldt. edu/ ~serc/ h2safety. html)". Humboldt State University. . Retrieved 2008-03-15.
Further reading • Chart of the Nuclides (http:/ / chartofthenuclides. com/ default. html). Fourteenth Edition. General Electric Company. 1989. http:/ / chartofthenuclides. com/ default. html. • Ferreira-Aparicio, P; M. J. Benito, J. L. Sanz (2005). "New Trends in Reforming Technologies: from Hydrogen Industrial Plants to Multifuel Microreformers". Catalysis Reviews 47: 491–588. doi: 10.1080/01614940500364958 (http:/ / dx. doi. org/ 10. 1080/ 01614940500364958). • Newton, David E. (1994). The Chemical Elements. New York, NY: Franklin Watts. ISBN 0-531-12501-7.
17
Hydrogen • Rigden, John S. (2002). Hydrogen: The Essential Element. Cambridge, MA: Harvard University Press. ISBN 0-531-12501-7. • Romm, Joseph, J. (2004). The Hype about Hydrogen, Fact and Fiction in the Race to Save the Climate. Island Press. ISBN 1-55963-703-X. Author interview (http:/ / www. globalpublicmedia. com/ transcripts/ 635) at Global Public Media. • Stwertka, Albert (2002). A Guide to the Elements. New York, NY: Oxford University Press. ISBN 0-19-515027-9.
External links • Basic Hydrogen Calculations of Quantum Mechanics (http:/ / www. physics. drexel. edu/ ~tim/ open/ hydrofin/ ) • Hydrogen phase diagram (http:/ / www. astro. washington. edu/ users/ larson/ Astro150b/ Lectures/ JupSatUraNep/ hydrogen_phase. gif) • Wavefunction of hydrogen (http:/ / hyperphysics. phy-astr. gsu. edu/ Hbase/ quantum/ hydwf. html#c3)
18
Article Sources and Contributors
Article Sources and Contributors Hydrogen Source: http://en.wikipedia.org/w/index.php?oldid=307051115 Contributors: (jarbarf), -- April, -jmac-, 0kdal, 16@r, 65.68.87.xxx, AStudent, Abarry, Acerty123, Acroterion, Adambro, Adashiel, Addshore, Adimovk5, Aervanath, Ahoerstemeier, Aitias, Akusarujin, Alexf, Alexfusco5, AlexiusHoratius, Allstarecho, Alrasheedan, Alsandro, Ancheta Wis, AndreasJS, Andrei Ramanonov, Andres, Andy M. Wang, Anonymous56789, Antandrus, Anthony Appleyard, Archimerged, Army1987, Arsonal, Ashmedai, Ashmoo, Atelaes, Atlant, AxelBoldt, AySz88, AzaToth, BAZZA42, BBird, BD2412, BRG, Basicdesign, Bcorr, Beetstra, BenFrantzDale, Bender235, Benjamnjoel2, Benji Franklyn, Bensaccount, Bibbidbabbidi, BigBen212, Blahbleh, Blainster, Blimpguy, Blind Man Walking, BlueEarth, Bobo192, Bongitybongbong, Bookandcoffee, BorgQueen, Borgdylan, Bpeps, BradBeattie, Brandonrush, Bravehart10000, Brendenhull, Brighterorange, Bryan Derksen, Buckyboy314, Bunnyhop11, Buster79, Bù hán ér lì, C'est moi, CStyle, CYD, CambridgeBayWeather, Cameron Nedland, Can't sleep, clown will eat me, Canageek, CaptainVindaloo, Carlobus, Carnildo, Casull, Celarnor, Cerealkiller13, Cfailde, ChaoticLlama, Charles Gaudette, Chemkid1, ChicXulub, Chochopk, Chowbok, Chris Dybala, Chrisjj, Chrislk02, Christopher Parham, Cimbalom, Civil Engineer III, Clicketyclack, Code E, Colbuckshot, Commander Keane, Condem, Contango, Conversion script, Coppro, Corporal butters, Cosmium, Costyn, Coviekiller5, Crescentnebula, Cryptic, Cryptic C62, CyclePat, Cyclotronwiki, Cyrius, D. 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Ufwuct, Ungvichian, Urhixidur, Uruiamme, Utcursch, Valencerian, Vanderdecken, Vegaswikian, Viddin66, Viskonsas, Vlectric, Vortexrealm, Vsmith, Vssun, Vuo, WAS 4.250, WJBscribe, WadeSimMiser, Waggers, Walkerma, Walton One, Warrenokeefe, Watch37264, Wavelength, Wayward, Webguy, WelshMatt, Whirling Sands, Wii Wiki, Wiki alf, Wikibob, Wikicali00, Wikipediarules2221, Wikispork, Wikiwert, William Allen Simpson, WillowW, Wimt, Wk muriithi, Wms editor, Wolfkeeper, Wtshymanski, Wyllium, XJamRastafire, Xaoa, Xenonice, Xevi, Xnuala, Xy7, Xyzzyplugh, Yath, Yidisheryid, Yinon, Yorktown1776, Zach4636, ZayZayEM, Zoe, 1504 anonymous edits
Image Sources, Licenses and Contributors file:hexagonal.svg Source: http://en.wikipedia.org/w/index.php?title=File:Hexagonal.svg License: BSD Contributors: Original uploader was Danieljamesscott at en.wikipedia file:Electron shell 001 Hydrogen.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_001_Hydrogen.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 Image:Shuttle Main Engine Test Firing cropped edited and reduced.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Shuttle_Main_Engine_Test_Firing_cropped_edited_and_reduced.jpg License: Public Domain Contributors: Avron, WTCA Image:hydrogen atom.svg Source: http://en.wikipedia.org/w/index.php?title=File:Hydrogen_atom.svg License: Public Domain Contributors: User:Bensaccount Image:Liquid hydrogen bubblechamber.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Liquid_hydrogen_bubblechamber.jpg License: Public Domain Contributors: Pieter Kuiper, Saperaud Image:Hydrogen.svg Source: http://en.wikipedia.org/w/index.php?title=File:Hydrogen.svg License: unknown Contributors: Mets501, Mion, Soeb, Treisijs, Xxxx00, 4 anonymous edits Image:Triangulum.nebula.full.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Triangulum.nebula.full.jpg License: unknown Contributors: Aarchiba, Dbenbenn, Denniss, Juiced lemon, Kluka, Lars Lindberg Christensen, Locos epraix, Tryphon, 5 anonymous edits
19
Image Sources, Licenses and Contributors Image:Emission spectrum-H.png Source: http://en.wikipedia.org/w/index.php?title=File:Emission_spectrum-H.png License: Public Domain Contributors: user:Merikanto Image:Main symptoms of hydrogen toxicity.png Source: http://en.wikipedia.org/w/index.php?title=File:Main_symptoms_of_hydrogen_toxicity.png License: Public Domain Contributors: Mikael Häggström
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
20
Helium
1
Helium hydrogen ← helium → lithiumNt ↑ He ↓ Ne
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
2He Periodic table
Appearance colorless gas General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensityMelting pointBoiling pointCritical pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure (defined by ITS-90) P/Pa
1
10
at T/K
100
1k
10 k
100 k
1.23
1.67
2.48
4.21
Atomic properties ElectronegativityIonization energies 2nd: 5250.5 kJ·mol−1Covalent radiusVan der Waals radius Miscellaneous Crystal structureMagnetic orderingThermal conductivityThermal expansionCAS registry number Most stable isotopes Main article: Isotopes of helium iso
N.A.
half-life
3
0.000137%*
3
4
99.999863%*
4
He He
DM
DE (MeV)
He is stable with 1 neutron He is stable with 2 neutron
*Atmospheric value, abundance may differ elsewhere.
DP
Helium helium, He, 2 noble gases 18, 1, s4.002602(2) g·mol−1 1s2 2 (Image) gas (0 °C, 101.325 kPa) 0.1786 g/L (at 2.5 MPa) 0.95 K,−272.20 °C,−457.96 °F 4.22 K,−268.93 °C,−452.07 °F 5.19 K, 0.227 MPa 0.0138 kJ·mol−1 0.0829 kJ·mol−1 (25 °C) 20.786 J·mol−1·K−1 no data (Pauling −1 [1] scale) 1st: 2372.3 kJ·mol 28 pm 140 pm hexagonal close-packed diamagnetic (300 K) −1 −1 −1 −1 0.1513 W·m ·K (25 °C) { µm·m ·K 7440-59-7 Helium (pronounced /ˈhiːliəm/) is the chemical element with atomic number 2, and is represented by the symbol He. It is a colorless, odorless, tasteless, non-toxic, inert monatomic gas that heads the noble gas group in the periodic table. Its boiling and melting points are the lowest among the elements and it exists only as a gas except in extreme conditions. An unknown yellow spectral line signature in sunlight was first observed from a solar eclipse in 1868 by French astronomer Pierre Janssen. Janssen is jointly credited with the discovery of the element with Norman Lockyer, who observed the same eclipse and was the first to propose that the line was due to a new element which he named helium. In 1903, large reserves of helium were found in the natural gas fields of the United States, which is by far the largest supplier of the gas. Helium is used in cryogenics, in deep-sea breathing systems, to cool superconducting magnets, in helium dating, for inflating balloons, for providing lift in airships and as a protective gas for many industrial uses (such as arc welding and growing silicon wafers). Inhaling a small volume of the gas temporarily changes the timbre and quality of the human voice. The behavior of liquid helium-4's two fluid phases, helium I and helium II, is important to researchers studying quantum mechanics (in particular the phenomenon of superfluidity) and to those looking at the effects that temperatures near absolute zero have on matter (such as superconductivity). Helium is the second lightest element and is the second most abundant in the observable universe, being present in the universe in masses more than 12 times those of all the other elements heavier than helium combined. Helium's abundance is also similar to this in our own Sun and Jupiter. This high abundance is due to the very high binding energy (per nucleon) of helium-4 with respect to the next three elements after helium (lithium, beryllium, and boron). This helium-4 binding energy also accounts for its commonality as a product in both nuclear fusion and radioactive decay. Most helium in the universe is helium-4, and was formed during the Big Bang. Some new helium is being created presently as a result of the nuclear fusion of hydrogen, in all but the very heaviest stars, which fuse helium into heavier elements at the extreme ends of their lives. On Earth, the lightness of helium has caused its evaporation from the gas and dust cloud from which the planet condensed, and it is thus relatively rare. What helium is present today has been mostly created by the natural radioactive decay of heavy radioactive elements (thorium and uranium), as the alpha particles that are emitted by such decays consist of helium-4 nuclei. This radiogenic helium is trapped with natural gas in concentrations up to seven percent by volume, from which it is extracted commercially by a low-temperature separation process called fractional distillation.
2
Helium
History Scientific discoveries The first evidence of helium was observed on August 18, 1868 as a bright yellow line with a wavelength of 587.49 nanometers in the spectrum of the chromosphere of the Sun. The line was detected by French astronomer Pierre Janssen during a total solar eclipse in Guntur, India.[2] [3] This line was initially assumed to be sodium. On October 20 of the same year, English astronomer Norman Lockyer observed a yellow line in the solar spectrum, which he named the D3 Fraunhofer line because it was near the known D1 and D2 lines of sodium.[4] He concluded that it was caused by an element in the Sun unknown on Earth. Lockyer and English chemist Edward Frankland named the element with the Greek word for the Sun, ἥλιος (helios).[5] [6] On March 26, 1895 British chemist Sir William Ramsay isolated helium on Earth by treating the mineral cleveite (a variety of uraninite with at least 10% rare earth elements) with mineral acids. Ramsay was looking for argon but, after separating Spectral lines of helium nitrogen and oxygen from the gas liberated by sulfuric acid, he noticed a bright yellow line that matched the D3 line observed in the spectrum of the Sun.[4] [7] [8] [9] These samples were identified as helium by Lockyer and British physicist William Crookes. It was independently isolated from cleveite in the same year by chemists Per Teodor Cleve and Abraham Langlet in Uppsala, Sweden, who collected enough of the gas to accurately determine its atomic weight.[3] [10] [11] Helium was also isolated by the American geochemist William Francis Hillebrand prior to Ramsay's discovery when he noticed unusual spectral lines while testing a sample of the mineral uraninite. Hillebrand, however, attributed the lines to nitrogen. His letter of congratulations to Ramsay offers an interesting case of discovery and near-discovery in science.[12] In 1907, Ernest Rutherford and Thomas Royds demonstrated that alpha particles are helium nuclei by allowing the particles to penetrate the thin glass wall of an evacuated tube, then creating a discharge in the tube to study the spectra of the new gas inside. In 1908, helium was first liquefied by Dutch physicist Heike Kamerlingh Onnes by cooling the gas to less than one kelvin.[13] He tried to solidify it by further reducing the temperature but failed because helium does not have a triple point temperature at which the solid, liquid, and gas phases are at equilibrium. Onnes' student Willem Hendrik Keesom was eventually able to solidify 1 cm3 of helium in 1926.[14] In 1938, Russian physicist Pyotr Leonidovich Kapitsa discovered that helium-4 has almost no viscosity at temperatures near absolute zero, a phenomenon now called superfluidity.[15] This phenomenon is related to Bose-Einstein condensation. In 1972, the same phenomenon was observed in helium-3, but at temperatures much closer to absolute zero, by American physicists Douglas D. Osheroff, David M. Lee, and Robert C. Richardson. The phenomenon in helium-3 is thought to be related to pairing of helium-3 fermions to make bosons, in analogy to Cooper pairs of electrons producing superconductivity.[16]
3
Helium
Extraction and use After an oil drilling operation in 1903 in Dexter, Kansas produced a gas geyser that would not burn, Kansas state geologist Erasmus Haworth collected samples of the escaping gas and took them back to the University of Kansas at Lawrence where, with the help of chemists Hamilton Cady and David McFarland, he discovered that the gas consisted of, by volume, 72% nitrogen, 15% methane (a combustible percentage only with sufficient oxygen), 1% hydrogen, and 12% an unidentifiable gas.[3] [17] With further analysis, Cady and McFarland discovered that 1.84% of the gas sample was helium.[18] [19] This showed that despite its overall rarity on Earth, helium was concentrated in large quantities under the American Great Plains, available for extraction as a byproduct of natural gas.[20] The greatest reserves of helium were in the Hugoton and nearby gas fields in southwest Kansas and the panhandles of Texas and Oklahoma. This enabled the United States to become the world's leading supplier of helium. Following a suggestion by Sir Richard Threlfall, the United States Navy sponsored three small experimental helium production plants during World War I. The goal was to supply barrage balloons with the non-flammable, lighter-than-air gas. A total of 200 thousand cubic feet (5,700 m3) of 92% helium was produced in the program even though only a few cubic feet (less than 100 liters) of the gas had previously been obtained.[4] Some of this gas was used in the world's first helium-filled airship, the U.S. Navy's C-7, which flew its maiden voyage from Hampton Roads, Virginia to Bolling Field in Washington, D.C. on December 1, 1921.[21] Although the extraction process, using low-temperature gas liquefaction, was not developed in time to be significant during World War I, production continued. Helium was primarily used as a lifting gas in lighter-than-air craft. This use increased demand during World War II, as well as demands for shielded arc welding. The helium mass spectrometer was also vital in the atomic bomb Manhattan Project.[22] The government of the United States set up the National Helium Reserve in 1925 at Amarillo, Texas with the goal of supplying military airships in time of war and commercial airships in peacetime.[4] Due to a US military embargo against Germany that restricted helium supplies, the Hindenburg was forced to use hydrogen as the lift gas. Helium use following World War II was depressed but the reserve was expanded in the 1950s to ensure a supply of liquid helium as a coolant to create oxygen/hydrogen rocket fuel (among other uses) during the Space Race and Cold War. Helium use in the United States in 1965 was more than eight times the peak wartime consumption.[23] After the "Helium Acts Amendments of 1960" (Public Law 86–777), the U.S. Bureau of Mines arranged for five private plants to recover helium from natural gas. For this helium conservation program, the Bureau built a 425 mile (684 km) pipeline from Bushton, Kansas to connect those plants with the government's partially depleted Cliffside gas field, near Amarillo, Texas. This helium-nitrogen mixture was injected and stored in the Cliffside gas field until needed, when it then was further purified.[24] By 1995, a billion cubic meters of the gas had been collected and the reserve was US$1.4 billion in debt, prompting the Congress of the United States in 1996 to phase out the reserve.[3] [25] The resulting "Helium Privatization Act of 1996"[26] (Public Law 104–273) directed the United States Department of the Interior to start emptying the reserve by 2005.[27]
4
Helium
5
Helium produced between 1930 and 1945 was about 98.3% pure (2% nitrogen), which was adequate for airships. In 1945, a small amount of 99.9% helium was produced for welding use. By 1949, commercial quantities of Grade A 99.95% helium were available.[28] For many years the United States produced over 90% of commercially usable helium in the world, while extraction plants in Canada, Poland, Russia, and other nations produced the remainder. In the mid-1990s, a new plant in Arzew, Algeria producing 600 million cubic feet (17 million cubic meters) began operation, with enough production to cover all of Europe's demand. Meanwhile, by 2000, the consumption of helium within the US had risen to above 15,000 metric tons.[29] In 2004–2006, two additional plants, one in Ras Laffen, Qatar and the other in Skikda, Algeria were built, but as of early 2007, Ras Laffen is functioning at 50%, and Skikda has yet to start up. Algeria quickly became the second leading producer of helium.[30] Through this time, both helium consumption and the costs of producing helium increased.[31] In the 2002 to 2007 period helium prices doubled,[32] and during 2008 alone the major suppliers raised prices about 50%.
Characteristics The helium atom Helium atom
An illustration of the helium atom, depicting the nucleus (pink) and the electron cloud distribution (black). The nucleus (upper right) in helium-4 is in reality spherically symmetric and closely resembles the electron cloud, although for more complicated nuclei this is not always the case. The black bar is one ångström, equal to 10−10 m or 100,000 fm.
Helium is the next simplest atom to solve using the rules of quantum mechanics, after the hydrogen atom. Helium is composed of two electrons in orbit around a nucleus containing two protons along with some neutrons. However, as in Newtonian mechanics, no system consisting of more than two particles can be solved with an exact analytical mathematical approach (see 3-body problem) and helium is no exception. Thus, numerical mathematical methods are required, even to solve the system of one nucleus and two electrons. The hydrogen atom quantum model has been used extensively to aid in solving the helium atom. The Niels Bohr model of the atom gave a very accurate explanation of the hydrogen
Helium spectrum, but when it came to helium, it collapsed. Werner Heisenberg developed a modification of Bohr's analysis but it involved half-integral values for the quantum numbers[33] . Thomas-Fermi theory also known as density functional theory is used to obtain the ground state energy levels of the helium atom along with the Hartree-Fock method. These methods have been used to create a quantum mechanical picture of helium electron binding which is accurate to within < 2% of the correct value, within a few numerical approximation steps. In such a model, various influences must be taken account of, including the electric repulsion of the electrons for each other, and the fact that one electron will, in part, screen the charge of the nucleus for the other. In the case of helium, it has been found that the effective nuclear charge "Z" which each electron sees, is about 1.69 units, not the 2 charges of a classic "bare" helium nucleus. The nucleus of the helium-4 atom, which is identical with an alpha particle is particularly interesting, inasmuch as high energy electron-scattering experiments show its charge to decrease exponentially from a maximum at a central point, exactly as does the charge density of helium's own electron cloud. The reason for this symmetry is elegant: the pair of neutrons and pair of protons in helium's nucleus both obey exactly the same quantum mechanical rules as do helium's pair of electrons (although the nuclear particles are subject to a different nuclear binding potential), so that all these fermions fully occupy 1s orbitals in pairs, none of them possessing orbital angular momentum, and each cancelling the other's intrinsic spin. This arrangement is energetically extremely stable for all these particles, and this stability accounts for many crucial facts regarding helium in nature. For example, the stability and low energy of the electron cloud state in helium accounts for the element's chemical inertness (the most extreme of all the elements), and also the lack of interaction of helium atoms with themselves, producing the lowest melting and boiling points of all the elements. In a similar way, the particular energetic stability of the helium-4 nucleus, produced by similar effects, accounts for the ease of helium-4 production in atomic reactions involving both heavy-particle emission, and fusion. The stability of helium-4 is the reason hydrogen is converted to helium-4 (not deuterium or helium-3 or heavier elements) in the Sun. It is also responsible for the fact the alpha particle is by far the most common type of baryonic particle to be ejected from atomic nuclei—that is, (alpha decay is far more common than cluster decay). The unusual stability of the helium-4 nucleus is also important cosmologically—it explains the fact that in the first few mintutes after the Big Bang, as the soup of free protons and neutrons which had been created in about 6:1 ratio, cooled to the point that nuclear binding was possible, the first nuclei to form were helium-4 nuclei. So tight was helium-4 binding, in fact, than it consumed nearly all of the free neutrons before they could beta-decay, leaving very few left to form any lithium, beryllium, or boron. Helium-4 nuclear binding is stronger than in any of these elements (see nucleogenesis and binding energy) and thus no energetic drive was available, once helium had been formed, to make elements 3, 4 and 5. It was barely energetically favorable for helium to fuse into the next element with a lower energy per nucleon, carbon. However, due to lack of intermediate elements, this process would take three helium nuclei striking each other nearly simultaneously (see triple alpha process). There was thus no time for significant carbon to be formed in the Big Bang, before the early expanding universe cooled in a matter of minutes to the temperature and pressure point where helium fusion to carbon was no longer possible. This left the early
6
Helium
7
universe with a very similar ratio of hydrogen to helium as is seen today (3 parts hydrogen to 1 part helium-4 by mass), with nearly all the neutrons in the universe (even as it exists today) trapped in the helium-4. All heavier elements (including those necessary for rocky planets like the Earth, and for carbon-based or other life), have thus had to be created since the Big Bang, in stars which were hot enough to burn not just hydrogen (for this produces only more helium), but hot enough to burn helium itself. Such stars are massive and therefore rare, and this fact accounts for the fact that all other chemical elements after hydrogen and helium today account for only 2% of the mass of atomic mater in the universe. Helium-4, by contrast, makes up about 23% of the universe's ordinary matter—nearly all the ordinary matter which isn't hydrogen.
Gas and plasma phases Helium is the least reactive noble gas after neon and thus the second least reactive of all elements; it is inert and monatomic in all standard conditions. Due to helium's relatively low molar (atomic) mass, in the gas phase its thermal conductivity, specific heat, and sound speed are all greater than any other gas except hydrogen. For similar reasons, and also due to the small size of helium atoms, helium's diffusion rate through solids is three times that of air and around 65% that of hydrogen.[4] Helium is less water soluble than any other gas
Helium discharge tube shaped like the element's atomic symbol
known,[34] and helium's index of refraction is closer to unity than that of any other gas.[35] Helium has a negative Joule-Thomson coefficient at normal ambient temperatures, meaning it heats up when allowed to freely expand. Only below its Joule-Thomson inversion temperature (of about 32 to 50 K at 1 atmosphere) does it cool upon free expansion.[4] Once precooled below this temperature, helium can be liquefied through expansion cooling.
Most extraterrestrial helium is found in a plasma state, with properties quite different from those of atomic helium. In a plasma, helium's electrons are not bound to its nucleus, resulting in very high electrical conductivity, even when the gas is only partially ionized. The charged particles are highly influenced by magnetic and electric fields. For example, in the solar wind together with ionized hydrogen, the particles interact with the Earth's magnetosphere giving rise to Birkeland currents and the aurora.[36]
Solid and liquid phases Unlike any other element, helium will remain liquid down to absolute zero at normal pressures. This is a direct effect of quantum mechanics: specifically, the zero point energy of the system is too high to allow freezing. Solid helium requires a temperature of 1–1.5 K (about −272 °C or −457 °F) and about 25 bar (2.5 MPa) of pressure.[37] It is often hard to distinguish solid from liquid helium since the refractive index of the two phases are nearly the same. The solid has a sharp melting point and has a crystalline structure, but it is highly compressible; applying pressure in a laboratory can decrease its volume by more than 30%.[38] With a bulk modulus on the order of 5×107 Pa[39] it is 50 times more compressible
Helium than water. Solid helium has a density of 0.214 ± 0.006 g/ml at 1.15 K and 66 atm; the projected density at 0 K and 25 bar is 0.187 ± 0.009 g/ml.[40] Helium I state Below its boiling point of 4.22 kelvin and above the lambda point of 2.1768 kelvin, the isotope helium-4 exists in a normal colorless liquid state, called helium I.[4] Like other cryogenic liquids, helium I boils when it is heated and contracts when its temperature is lowered. Below the lambda point, however, helium doesn't boil, and it expands as the temperature is lowered further. Helium I has a gas-like index of refraction of 1.026 which makes its surface so hard to see that floats of styrofoam are often used to show where the surface is.[4] This colorless liquid has a very low viscosity and a density one-eighth that of water, which is only one-fourth the value expected from classical physics.[4] Quantum mechanics is needed to explain this property and thus both types of liquid helium are called quantum fluids, meaning they display atomic properties on a macroscopic scale. This may be an effect of its boiling point being so close to absolute zero, preventing random molecular motion (thermal energy) from masking the atomic properties.[4] Helium II state Liquid helium below its lambda point begins to exhibit very unusual characteristics, in a state called helium II. Boiling of helium II is not possible due to its high thermal conductivity; heat input instead causes evaporation of the liquid directly to gas. The isotope helium-3 also has a superfluid phase, but only at much lower temperatures; as a result, less is known about such properties in the isotope helium-3.[4]
8
Helium
9
Helium II is a superfluid, a quantum-mechanical state of matter with strange properties. For example, when it flows through capillaries as thin as 10−7 to 10−8 m it has no measurable viscosity.[3] However, when measurements were done between two moving discs, a viscosity comparable to that of gaseous helium was observed. Current theory explains this using the two-fluid model for helium II. In this model, liquid helium below the lambda point is viewed as containing a proportion of helium atoms in a ground state, which are superfluid and flow with exactly zero viscosity, and a proportion of helium atoms in an excited state, which behave more like an ordinary fluid.[41] In the fountain effect, a chamber is constructed which is connected to a reservoir of helium II by a sintered disc through which superfluid helium leaks easily but through which non-superfluid helium cannot pass. If the interior of the container is heated, the superfluid helium changes to non-superfluid helium. In order to maintain the equilibrium fraction of superfluid helium, superfluid helium leaks through and increases the pressure, causing liquid to fountain out of the container.[42]
Unlike ordinary liquids, helium II will creep along surfaces in order to reach an equal level; after a short while, the levels in the two containers will equalize. The Rollin film also covers the interior of the larger container; if it were not sealed, the [4] helium II would creep out and escape.
The thermal conductivity of helium II is greater than that of any other known substance, a million times that of helium I and several hundred times that of copper.[4] This is because heat conduction occurs by an exceptional quantum-mechanical mechanism. Most materials that conduct heat well have a valence band of free electrons which serve to transfer the heat. Helium II has no such valence band but nevertheless conducts heat well. The flow of heat is governed by equations that are similar to the wave equation used to characterize sound propagation in air. When heat is introduced, it moves at 20 meters per second at 1.8 K through helium II as waves in a phenomenon known as second sound.[4] Helium II also exhibits a creeping effect. When a surface extends past the level of helium II, the helium II moves along the surface, seemingly against the force of gravity. Helium II will escape from a vessel that is not sealed by creeping along the sides until it reaches a warmer region where it evaporates. It moves in a 30 nm-thick film regardless of surface material. This film is called a Rollin film and is named after the man who first characterized this trait, Bernard V. Rollin.[4] [43] [44] As a result of this creeping behavior and helium II's ability to leak rapidly through tiny openings, it is very difficult to confine liquid helium. Unless the container is carefully constructed, the helium II will creep along the surfaces and through valves until it reaches somewhere warmer, where it will evaporate. Waves propagating across a Rollin film are governed by the same equation as gravity waves in shallow water, but rather than gravity, the restoring force is the Van der Waals force.[45] These waves are known as third sound.[46]
Helium
Isotopes There are eight known isotopes of helium, but only helium-3 and helium-4 are stable. In the Earth's atmosphere, there is one 3He atom for every million 4He atoms.[3] Unlike most elements, helium's isotopic abundance varies greatly by origin, due to the different formation processes. The most common isotope, helium-4, is produced on Earth by alpha decay of heavier radioactive elements; the alpha particles that emerge are fully ionized helium-4 nuclei. Helium-4 is an unusually stable nucleus because its nucleons are arranged into complete shells. It was also formed in enormous quantities during Big Bang nucleosynthesis.[47] Helium-3 is present on Earth only in trace amounts; most of it since Earth's formation, though some falls to Earth trapped in cosmic dust.[48] Trace amounts are also produced by the beta decay of tritium.[49] Rocks from the Earth's crust have isotope ratios varying by as much as a factor of ten, and these ratios can be used to investigate the origin of rocks and the composition of the Earth's mantle.[48] 3He is much more abundant in stars, as a product of nuclear fusion. Thus in the interstellar medium, the proportion of 3He to 4He is around 100 times higher than on Earth.[50] Extraplanetary material, such as lunar and asteroid regolith, have trace amounts of helium-3 from being bombarded by solar winds. The Moon's surface contains helium-3 at concentrations on the order of 0.01 ppm.[51] [52] A number of people, starting with Gerald Kulcinski in 1986,[53] have proposed to explore the moon, mine lunar regolith and use the helium-3 for fusion. Liquid helium-4 can be cooled to about 1 kelvin using evaporative cooling in a 1-K pot. Similar cooling of helium-3, which has a lower boiling point, can achieve about 0.2 kelvin in a helium-3 refrigerator. Equal mixtures of liquid 3He and 4He below 0.8 K separate into two immiscible phases due to their dissimilarity (they follow different quantum statistics: helium-4 atoms are bosons while helium-3 atoms are fermions).[4] Dilution refrigerators use this immiscibility to achieve temperatures of a few millikelvins. It is possible to produce exotic helium isotopes, which rapidly decay into other substances. The shortest-lived heavy helium isotope is helium-5 with a half-life of 7.6 × 10−22 seconds. Helium-6 decays by emitting a beta particle and has a half life of 0.8 seconds. Helium-7 also emits a beta particle as well as a gamma ray. Helium-7 and helium-8 are created in certain nuclear reactions.[4] Helium-6 and helium-8 are known to exhibit a nuclear halo. Helium-2 (two protons, no neutrons) is a radioisotope that decays by proton emission into protium, with a half-life of 3 × 10−27 seconds.[4]
Compounds Helium is chemically unreactive under all normal conditions due to its valence of zero.[38] It is an electrical insulator unless ionized. As with the other noble gases, helium has metastable energy levels that allow it to remain ionized in an electrical discharge with a voltage below its ionization potential.[4] Helium can form unstable compounds, known as excimers, with tungsten, iodine, fluorine, sulfur and phosphorus when it is subjected to an electric glow discharge, to electron bombardment, or else is a plasma for another reason. HeNe, HgHe10, WHe2 and the molecular ions He+2, He2+2, HeH+, and HeD+ have been created this way.[54] This technique has also allowed the production of the neutral molecule He2, which has a large number of band systems, and HgHe, which is apparently only held together by polarization forces.[4] Theoretically, other true compounds may also be possible, such as helium fluorohydride (HHeF) which would be analogous to HArF,
10
Helium discovered in 2000.[55] . Calculations show that two new compounds containing a helium-oxygen bond could be stable.[56] . The two new molecular species, predicted using theory, CsFHeO and N(CH3)4FHeO, are derivatives of a metastable [F– HeO] anion first theorized in 2005 by a group from Taiwan. If confirmed by experiment such compounds will [57] end helium's chemical nobility, and the only remaining noble element will be neon. Helium has been put inside the hollow carbon cage molecules (the fullerenes) by heating under high pressure. The endohedral fullerene molecules formed are stable up to high temperatures. When chemical derivatives of these fullerenes are formed, the helium stays inside.[58] If helium-3 is used, it can be readily observed by helium nuclear magnetic resonance spectroscopy.[59] Many fullerenes containing helium-3 have been reported. Although the helium atoms are not attached by covalent or ionic bonds, these substances have distinct properties and a definite composition, like all stoichiometric chemical compounds.
Occurrence and production Natural abundance Helium is the second most abundant element in the known Universe (after hydrogen), constituting 23% of the baryonic mass of the Universe.[3] The vast majority of helium was formed by Big Bang nucleosynthesis from one to three minutes after the Big Bang. As such, measurements of its abundance contribute to cosmological models. In stars, it is formed by the nuclear fusion of hydrogen in proton-proton chain reactions and the CNO cycle, part of stellar nucleosynthesis.[47] In the Earth's atmosphere, the concentration of helium by volume is only 5.2 parts per million.[60] [61] The concentration is low and fairly constant despite the continuous production of new helium because most helium in the Earth's atmosphere escapes into space by several processes.[62] [63] In the Earth's heterosphere, a part of the upper atmosphere, helium and other lighter gases are the most abundant elements. Nearly all helium on Earth is a result of radioactive decay, and thus an Earthly helium balloon is essentially a bag of retired alpha particles. Helium is found in large amounts in minerals of uranium and thorium, including cleveites, pitchblende, carnotite and monazite, because they emit alpha particles (helium nuclei, He2+) to which electrons immediately combine as soon as the particle is stopped by the rock. In this way an estimated 3000 tonnes of helium are generated per year throughout the lithosphere.[64] [65] [66] In the Earth's crust, the concentration of helium is 8 parts per billion. In seawater, the concentration is only 4 parts per trillion. There are also small amounts in mineral springs, volcanic gas, and meteoric iron. Because helium is trapped in a similar way by non-permeable layer of rock like natural gas the greatest concentrations on the planet are found in natural gas, from which most commercial helium is derived. The concentration varies in a broad range from a few ppm up to over 7% in a small gas field in San Juan County, New Mexico.[67] [68]
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Modern extraction For large-scale use, helium is extracted by fractional distillation from natural gas, which contains up to 7% helium.[69] Since helium has a lower boiling point than any other element, low temperature and high pressure are used to liquefy nearly all the other gases (mostly nitrogen and methane). The resulting crude helium gas is purified by successive exposures to lowering temperatures, in which almost all of the remaining nitrogen and other gases are precipitated out of the gaseous mixture. Activated charcoal is used as a final purification step, usually resulting in 99.995% pure Grade-A helium.[4] The principal impurity in Grade-A helium is neon. In a final production step, most of the helium that is produced is liquefied via a cryogenic process. This is necessary for applications requiring liquid helium and also allows helium suppliers to reduce the cost of long distance transportation, as the largest liquid helium containers have more than five times the capacity of the largest gaseous helium tube trailers.[30] [70] In 2005, approximately 160 million cubic meters of helium were extracted from natural gas or withdrawn from helium reserves, with approximately 83% from the United States, 11% from Algeria, and most of the remainder from Russia and Poland.[71] In the United States, most helium is extracted from natural gas of the Hugoton and nearby gas fields in Kansas, Oklahoma, and Texas.[30] Diffusion of crude natural gas through special semipermeable membranes and other barriers is another method to recover and purify helium.[72] Helium can be synthesized by bombardment of lithium or boron with high-velocity protons, but this is not an economically viable method of production.[73]
Applications Helium is used for many purposes that require some of its unique properties, such as its low boiling point, low density, low solubility, high thermal conductivity, or inertness. Helium is commercially available in either liquid or gaseous form. As a liquid, it can be supplied in small containers called Dewars which hold up to 1,000 liters of helium, or in large ISO containers which have nominal capacities as large as 11,000 US gallons (42 m3). In gaseous form, small quantities of helium are supplied in high pressure cylinders holding up to 300 standard cubic feet, while large quantities of high pressure gas are supplied in tube trailers which have capacities of up to 180,000 standard cubic feet. Airships, balloons and rocketry Because it is lighter than air, airships and balloons are inflated with helium for lift. While hydrogen gas is approximately 7% more buoyant, helium has the advantage of being non-flammable (in addition to being Because of its low density and fire retardant).[25] In rocketry, helium is used as an incombustibility, helium is the gas of ullage medium to displace fuel and oxidizers in storage choice to fill airships such as the tanks and to condense hydrogen and oxygen to make Goodyear blimp. rocket fuel. It is also used to purge fuel and oxidizer from ground support equipment prior to launch and to pre-cool liquid hydrogen in space vehicles. For example, the Saturn V booster used in the Apollo program needed about 13 million cubic feet (370,000 m3) of helium to launch.[38] Commercial and recreational
Helium
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Helium alone is less dense than atmospheric air, so it will change the timbre (not pitch[74] ) of a person's voice when inhaled. However, inhaling it from a typical commercial source, such as that used to fill balloons, can be dangerous due to the risk of asphyxiation from lack of oxygen, and the number of contaminants that may be present. These could include trace amounts of other gases, in addition to aerosolized lubricating oil. For its low solubility in nervous tissue, helium mixtures such as trimix, heliox and heliair are used for deep diving to reduce the effects of narcosis.[75] [76] At depths below 150 metres (490 ft) small amounts of hydrogen are added to a helium-oxygen mixture to counter the effects of high pressure nervous syndrome.[77] At these depths the low density of helium is found to considerably reduce the effort of breathing.[78] Helium-neon lasers have various applications, including barcode readers.[3] Industrial For its inertness and high thermal conductivity, neutron transparency, and because it does not form radioactive isotopes under reactor conditions, helium is used as a heat-transfer medium in some gas-cooled nuclear reactors.[79] Helium is used as a shielding gas in arc welding processes on materials that are contaminated easily by air.[3] Helium is used as a protective gas in growing silicon and germanium crystals, in titanium and zirconium production, and in gas chromatography,[38] because it is inert. Because of its inertness, thermally and calorically perfect nature, high speed of sound, and high value of the heat capacity ratio, it is also useful in supersonic wind tunnels[80] and impulse facilities[81] . Because it diffuses through solids at three times the rate of air, helium is used as a tracer gas to detect leaks in high-vacuum equipment and high-pressure containers.[79] Helium, mixed with a heavier gas such as xenon, is useful for thermoacoustic refrigeration due to the resulting high heat capacity ratio and low Prandtl number.[82] The inertness of helium has environmental advantages over conventional refrigeration systems which contribute to ozone depletion or global warming.[83] Scientific The use of helium reduces the distorting effects of temperature variations in the space between lenses in some telescopes, due to its extremely low index of refraction.[4] This method is especially used in solar telescopes where a vacuum tight telescope tube would be too heavy.[84] [85] The age of rocks and minerals that contain uranium and thorium can be estimated by measuring the level of helium with a process known as helium dating.[3] [4]
Liquid helium is used to cool the superconducting magnets in modern MRI scanners.
Liquid helium is used to cool certain metals to the extremely low temperatures required for superconductivity, such as in superconducting magnets for magnetic resonance imaging. The Large Hadron Collider at CERN uses 96 tonnes of liquid helium to maintain the temperature at 1.9 Kelvin.[86] Helium at low temperatures is also used in cryogenics. Helium is a commonly used carrier gas for gas chromatography. The leak rate of industrial vessels (typically vacuum chambers and cryogenic tanks) is measured using helium because of its small molecular diameter and because it is inert. No other inert substance will leak
Helium through micro-cracks or micro-pores in a vessel's wall at a greater rate than helium. A helium leak detector (see Helium mass spectrometer) is used to find leaks in vessels. Helium leaks through cracks should not be confused with gas permeation through a bulk material. While helium has documented permeation constants (thus a calculable permeation rate) through glasses, ceramics, and syntheic materials, inert gasses such as helium will not permeate most bulk metals.[87]
Safety Neutral helium at standard conditions is non-toxic, plays no biological role and is found in trace amounts in human blood. If enough helium is inhaled that oxygen needed for normal respiration is replaced asphyxia is possible. The safety issues for cryogenic helium are similar to those of liquid nitrogen; its extremely low temperatures can result in cold burns and the liquid to gas expansion ratio can cause explosions if no pressure-relief devices are installed. Containers of helium gas at 5 to 10 K should be handled as if they contain liquid helium due to the rapid and significant thermal expansion that occurs when helium gas at less than 10 K is warmed to room temperature.[38]
Biological effects The human voice is not like a string instrument, in which the a primarily vibrating object completely sets the pitch of the sound. Rather, in a human, the vocal folds act as a source of polytonic vibration, much like the reed(s) in woodwind musical instruments. As in a woodwind, the size of the resonant cavity plays a large part in picking out and amplifying a given fundamental or overtone frequency of vibration, during soundmaking. The voice of a person who has inhaled helium temporarily changes in timbre in a way that makes it sound high-pitched, because higher overtones are being amplified. The speed of sound in helium is nearly three times the speed of sound in air; because the fundamental frequency of a gas-filled cavity is proportional to the speed of sound in the gas, when helium is inhaled there is a corresponding increase in the pitch of the resonant frequencies of the vocal tract.[3] [88] (The opposite effect, lowering frequencies, can be obtained by inhaling a dense gas such as sulfur hexafluoride.) Inhaling helium can be dangerous if done to excess, since helium is a simple asphyxiant and so displaces oxygen needed for normal respiration.[3] [89] Breathing pure helium continuously causes death by asphyxiation within minutes. Inhaling helium directly from pressurized cylinders is extremely dangerous, as the high flow rate can result in barotrauma, fatally rupturing lung tissue.[89] [90] However, death caused by helium is quite rare, with only two fatalities reported between 2000 and 2004 in the United States.[90] At high pressures (more than about 20 atm or two MPa), a mixture of helium and oxygen (heliox) can lead to high pressure nervous syndrome, a sort of reverse-anesthetic effect; adding a small amount of nitrogen to the mixture can alleviate the problem.[91] [92]
14
Helium
15
See also • • • • • • •
Abiogenic petroleum origin Helium-3 propulsion Leidenfrost effect Quantum solid Superfluid Tracer-gas leak testing method Helium atom
References • Bureau of Mines (1967). Minerals yearbook mineral fuels Year 1965, Volume II (1967). U. S. Government Printing Office. • "Chart of the Nuclides: Fourteenth Edition [93]". General Electric Company. 1989. http:/ / chartofthenuclides. com/ default. html. • Emsley, John (1998). The Elements (3rd ed.). New York: Oxford University Press. ISBN 978-0198558187. • "Mineral Information for Helium [94]" (PDF). United States Geological Survey (usgs.gov). http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ helium/ heliumcs07. pdf. Retrieved on 2007-01-05. • Vercheval, J. (January 2003). "The thermosphere: a part of the heterosphere [95]". Belgian Institute for Space Aeronomy. http:/ / web. archive. org/ web/ 20050101090349/ www. oma. be/ BIRA-IASB/ Public/ Research/ Thermo/ Thermotxt. en. html. Retrieved on 2008-07-12. • Zastenker, G. N.; E. Salerno, F. Buehler, P. Bochsler, M. Bassi, Y. N. Agafonov, N. A. Eismont, V. V. Khrapchenkov, H. Busemann (April 2002). "Isotopic Composition and Abundance of Interstellar Neutral Helium Based on Direct Measurements [96]". Astrophysics 45 (2): 131–142. doi:10.1023/A:1016057812964 [97]. http:/ / www. ingentaconnect. com/ content/ klu/ asys/ 2002/ 00000045/ 00000002/ 00378626.
External links General • The Periodic Table of Videos - Helium
[98]
• US Government' Bureau of Land Management: Sources, Refinement, and Shortage. With some History of Helium. • U.S. Geological Survey Publicationson Helium [100] beginning 1996 • It's Elemental – Helium [101]
[99]
More detail • Helium [102] at the Helsinki University of Technology; includes pressure-temperature phase diagrams for helium-3 and helium-4 • Lancaster University, Ultra Low Temperature Physics [103] - includes a summary of some low temperature techniques Miscellaneous • Physics in Speech [104] with audio samples that demonstrate the unchanged voice pitch • Article about helium and other noble gases [105]
Helium
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16
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ISSN 0093-5387 (http:/ / worldcat. org/ issn/ 0093-5387). OCLC 2068005 (http:/ / worldcat. org/ oclc/ 2068005). PMID 3212843. . Retrieved on 2008-06-24. [78] Butcher, Scott J.; Richard L. Jones, Jonathan R. Mayne, Timothy C. Hartley, Stewart R. Petersen (December 2007). "Impaired exercise ventilatory mechanics with the self-contained breathing apparatus are improved with heliox". European Journal of Applied Physiology (Netherlands: Springer) 101 (6): 659(11). doi: 10.1007/s00421-007-0541-5 (http:/ / dx. doi. org/ 10. 1007/ s00421-007-0541-5). [79] Considine, Glenn D., ed (2005). "Helium". Van Nostrand's Encyclopedia of Chemistry. Wylie-Interscience. pp. 764–765. ISBN 0-471-61525-0. [80] Beckwith, I.E.; C. G. Miller III (1990). "Aerothermodynamics and Transition in High-Speed Wind Tunnels at Nasa Langley". Annual Review of Fluid Mechanics 22: 419–439. doi: 10.1146/annurev.fl.22.010190.002223 (http:/ / dx. doi. org/ 10. 1146/ annurev. fl. 22. 010190. 002223). [81] Morris, C.I. (2001) (PDF). Shock Induced Combustion in High Speed Wedge Flows (http:/ / thermosciences. stanford. edu/ pdf/ TSD-143. pdf). Stanford University Thesis. . [82] Belcher, James R.; William V. Slaton, Richard Raspet, Henry E. Bass, Jay Lightfoot (1999). "Working gases in thermoacoustic engines". The Journal of the Acoustical Society of America 105 (5): 2677–2684. doi: 10.1121/1.426884 (http:/ / dx. doi. org/ 10. 1121/ 1. 426884). [83] Makhijani, Arjun; Kevin Gurney (1995). Mending the Ozone Hole: Science, Technology, and Policy. MIT Press. ISBN 0262133083. [84] Jakobsson, H. (1997). "Simulations of the dynamics of the Large Earth-based Solar Telescope". Astronomical & Astrophysical Transactions 13 (1): 35–46. doi: 10.1080/10556799708208113 (http:/ / dx. doi. org/ 10. 1080/ 10556799708208113). [85] Engvold, O.; R.B. Dunn, R. N. Smartt, W. C. Livingston (1983). " Tests of vacuum VS helium in a solar telescope (http:/ / adsabs. harvard. edu/ cgi-bin/ nph-bib_query?bibcode=1983ApOpt. . 22. . . 10E& amp;db_key=AST)". Applied Optics 22: 10–12. doi: 10.1364/AO.22.000010 (http:/ / dx. doi. org/ 10. 1364/ AO. 22. 000010). . Retrieved on 2008-07-27. [86] " LHC Guide booklet CERN - LHC: Facts and Figures (http:/ / visits. web. cern. ch/ visits/ guides/ tools/ presentation/ LHC_booklet-2. pdf)". CERN. LHC Guide booklet. Retrieved on 2008-04-30. [87] Jack W. Ekin (2006). Experimental Techniques for Low-Temperature measurements (http:/ / books. google. co. jp/ books?id=Q9tmZQTDPiYC). Oxford University Press. ISBN 0198570546. . [88] Ackerman MJ, Maitland G (December 1975). " Calculation of the relative speed of sound in a gas mixture (http:/ / archive. rubicon-foundation. org/ 2738)". Undersea Biomed Res 2 (4): 305–10. PMID 1226588. . Retrieved on 2008-08-09. [89] () Grassberger, Martin; Astrid Krauskopf (2007). "Suicidal asphyxiation with helium: Report of three cases Suizid mit Helium Gas: Bericht über drei Fälle" (in German & English). Wiener Klinische Wochenschrift 119 (9–10): 323–325. doi: 10.1007/s00508-007-0785-4 (http:/ / dx. doi. org/ 10. 1007/ s00508-007-0785-4). [90] Engber, Daniel (2006-06-13). " Stay Out of That Balloon! (http:/ / www. slate. com/ id/ 2143631/ )". Slate.com. . Retrieved on 2008-07-14. [91] Rostain JC, Lemaire C, Gardette-Chauffour MC, Doucet J, Naquet R (April 1983). " Estimation of human susceptibility to the high-pressure nervous syndrome (http:/ / jap. physiology. org/ cgi/ pmidlookup?view=long&
19
Helium pmid=6853282)". J Appl Physiol 54 (4): 1063–70. PMID 6853282. . Retrieved on 2008-08-09. [92] Hunger Jr, W. L.; P. B. Bennett. (1974). " The causes, mechanisms and prevention of the high pressure nervous syndrome (http:/ / archive. rubicon-foundation. org/ 2661)". Undersea Biomed. Res. 1 (1): 1–28. ISSN 0093-5387 (http:/ / worldcat. org/ issn/ 0093-5387). OCLC 2068005 (http:/ / worldcat. org/ oclc/ 2068005). PMID 4619860. . Retrieved on 2008-08-09. [93] http:/ / chartofthenuclides. com/ default. html [94] http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ helium/ heliumcs07. pdf [95] http:/ / web. archive. org/ web/ 20050101090349/ www. oma. be/ BIRA-IASB/ Public/ Research/ Thermo/ Thermotxt. en. html [96] http:/ / www. ingentaconnect. com/ content/ klu/ asys/ 2002/ 00000045/ 00000002/ 00378626 [97] http:/ / dx. doi. org/ 10. 1023%2FA%3A1016057812964 [98] http:/ / uk. youtube. com/ watch?v=a8FJEiI5e6Q [99] http:/ / www. blm. gov/ wo/ st/ en/ info/ newsroom/ 2007/ january/ NR0701_2. html [100] http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ helium/ [101] http:/ / education. jlab. org/ itselemental/ ele002. html [102] http:/ / boojum. hut. fi/ research/ theory/ helium. html [103] http:/ / www. lancs. ac. uk/ depts/ physics/ research/ condmatt/ ult/ index. html [104] http:/ / www. phys. unsw. edu. au/ PHYSICS_!/ SPEECH_HELIUM/ speech. html [105] http:/ / www. du. edu/ ~jcalvert/ phys/ helium. htm
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21
Lithium
1
Lithium helium ← lithium → berylliumH ↑ Li ↓ Na
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
3Li Periodic table
Appearance silvery white (seen here in oil)
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointCritical pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
797
885
995
1144
1337
1610
Atomic properties Oxidation states ElectronegativityIonization energies 2nd: 7298.1 kJ·mol−1 3rd: 11815.0 kJ·mol−1Atomic radiusCovalent radiusVan der Waals radius
Lithium
2
Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusMohs hardnessCAS registry number Most stable isotopes Main article: Isotopes of lithium iso
N.A.
half-life
6
7.5%
6
7
92.5%
7
Li Li
DM
DE (MeV)
DP
Li is stable with 3 neutron Li is stable with 4 neutron 6
Li content may be as low as 3.75% in natural samples. 7Li would therefore have a content of up to 96.25%.
lithium, Li, 3 alkali metal1, 2, s6.941(2) g·mol−1 1s2 2s1 2, 1 (Image) solid 0.534 g·cm−3 0.512 g·cm−3 453.69 K,180.54 °C,356.97 °F 1615 K,1342 °C,2456.6 °F (extrapolated) 3223 K, 67 MPa 3.00 kJ·mol−1 147.1 kJ·mol−1 (25 °C) 24.860 J·mol−1·K−1+1, -1 (strongly basic oxide) 0.98 (Pauling scale) 1st: 520.2 kJ·mol−1152 pm128±7 pm 182 pm body-centered cubic paramagnetic (20 °C) 92.8 nΩ·m (300 K) 84.8 W·m−1·K−1 (25 °C) 46 µm·m−1·K−1 (20 °C) 6000 m/s 4.9 GPa 4.2 GPa 11 GPa 0.6 7439-93-2 Lithium (pronounced /ˈlɪθiəm/) is a soft, silver-white metal that belongs to the alkali metal group of chemical elements. It is represented by the symbol Li and has an atomic number of 3. Under standard conditions it is the lightest metal and the least dense solid element. Like all alkali metals lithium is highly reactive, corroding quickly in moist air to form a black tarnish. For this reason lithium metal is typically stored under the cover of oil. When cut open lithium exhibits a metallic luster, but contact with oxygen quickly turns it back to a dull silvery gray color. Lithium in its elemental state is highly flammable. According to theory, lithium was one of the few elements synthesized in the Big Bang. Since its current estimated abundance in the universe is vastly less than that predicted by theory[1] ; the processes by which new lithium is created and destroyed, and the true value of its abundance,[2] continue to be active matters of study in astronomy.[3] [4] [5] The nuclei of lithium are relatively fragile: the two stable lithium isotopes found in nature have lower binding energies per nucleon than any other stable compound nuclides, save for the exotic and rare deuterium, and 3He. [6] Though very light in atomic weight, lithium is less common in the solar system than 25 of the first 32 chemical elements. [7] Due to its high reactivity it only appears naturally in the form of compounds. Lithium occurs in a number of pegmatitic minerals, but is also commonly obtained from brines and clays. On a commercial scale, lithium metal is isolated electrolytically from a mixture of lithium chloride and potassium chloride. Trace amounts of lithium are present in the oceans and in some organisms, though the element serves no apparent vital biological function in humans. However, the lithium ion Li+ administered as any of several lithium salts has proved to be useful as a mood stabilizing drug due to neurological effects of the ion in the human body. Lithium and its compounds have several industrial applications, including heat-resistant glass and ceramics, high strength-to-weight alloys used in aircraft, and lithium batteries. Lithium also has important links to nuclear physics. The transmutation of lithium atoms to tritium was the first man-made form of a nuclear fusion reaction, and lithium deuteride serves as a fusion fuel in staged thermonuclear weapons.
Lithium
3
History and etymology Petalite (LiAlSi4O10, or lithium aluminum silicate) was first discovered in 1800 by the Brazilian scientist José Bonifácio de Andrade e Silva, who discovered the mineral in a Swedish mine on the island of Utö.[8] [9] [10] However, it was not until 1817 that Johan August Arfwedson, then working in the laboratory of Jöns Jakob Berzelius, discovered the presence of a new element while analyzing petalite ore.[11] [12] [13] The element formed compounds similar to those of sodium and potassium, though its carbonate and hydroxide were less water soluble and more basic.[14] Berzelius gave the alkaline material the name "lithos", from the Greek λιθoς (lithos, "stone"), to reflect its discovery in a mineral, as opposed to sodium and potassium, which had been discovered in plant tissue; its name was later standardized as "lithium".[13] [9] [15] Arfwedson later showed that this same element was present in the mineral ores spodumene and lepidolite. In 1818 Christian Gmelin was the first to observe that lithium salts give a bright red color in flame. However, both Arfwedson and Gmelin tried and failed to isolate the element from its salts.[16] [17] [13] The element was not isolated until 1821, when William Thomas Brande isolated the element by performing electrolysis on lithium oxide, a process previously employed by Sir Humphry Davy to isolate potassium and sodium.[17] [18] [19] Brande also described pure salts of lithium, such as the chloride, and performed an estimate of its atomic weight. In 1855 Robert Bunsen and Augustus Matthiessen produced large quantities of the metal by electrolysis of lithium chloride. The discovery of this procedure eventually led to commercial production of lithium metal, begun in 1923 by the German company Metallgesellschaft AG through the electrolysis of a molten mixture of lithium chloride and potassium chloride.[16] [20]
Properties Like the other alkali metals, lithium has a single valence electron that is easily given up to form a cation.[15] Because of this, it is both a good conductor of both heat and electricity and highly reactive, though it is the least reactive of the alkali metals due to the proximity of its valence electron to the nucleus.[15] Lithium is soft enough to be cut with a knife; it is the lightest and softest of the metals on the periodic table. When cut, it possesses a silvery-white color that quickly Lithium pellets (covered in white changes to gray due to oxidation.[15] It also has a low lithium hydroxide) density (approximately 0.534 g/cm3) and thus will float on water, though it reacts easily with water. This reaction is energetic, forming hydrogen gas and lithium hydroxide in aqueous solution.[15] Due to its reactivity with water, lithium is usually stored under cover of mineral oil or kerosene.[15] Lithium possesses a low coefficient of thermal expansion and the highest specific heat capacity of any solid element. Lithium is superconductive below 400 μK at standard pressure[21] and at higher temperatures (more than 9 kelvin) at very high pressures (over 200,000 atmospheres)[22] At cryogenic temperatures, lithium, like sodium, undergoes diffusionless phase change transformations. At 4.2K it has a rhombohedral crystal system (with a nine-layer repeat spacing)[23] ; at higher temperatures it transforms to
Lithium face-centered cubic and then body-centered cubic. At liquid-helium temperatures (4 K) the rhombohedral structure is the most prevalent.
Chemistry In moist air, lithium metal rapidly tarnishes to form a black coating of lithium hydroxide (LiOH and LiOH·H2O), lithium nitride (Li3N) and lithium carbonate (Li2CO3, the result of a secondary reaction between LiOH and CO2).[24] When placed over a flame lithium gives off a striking crimson color, but when it burns strongly the flame becomes a brilliant white. Lithium will ignite and burn in oxygen when exposed to water or water vapours. It is the only metal that reacts with nitrogen at room temperature. Lithium metal is flammable and potentially explosive when exposed to air and especially water, though less so than other alkali metals. The lithium-water reaction at normal temperatures is brisk but not violent, though the hydrogen produced can ignite. Like all alkali metals, lithium fires are difficult to extinguish, requiring dry powder fire extinguishers, specifically Class D type (see Types of extinguishing agents).
Lithium compounds Lithium has a diagonal relationship with magnesium, an element of similar atomic and ionic radius. Chemical resemblances between the two metals include the formation of a nitride by reaction with N2, the formation of an oxide when burnt in O2, salts with similar solubilities, and thermal instability of the carbonates and nitrides.[24]
Isotopes Naturally occurring lithium is composed of two stable isotopes, 6Li and 7Li, the latter being the more abundant (92.5 percent natural abundance).[25] [15] Both natural isotopes have anomalously low nuclear binding energy per nucleon compared to the next lighter and heavier elements, Helium and Beryllium, which means that alone among stable light elements, Lithium can produce net energy through nuclear fission. Seven radioisotopes have been characterized, the most stable being 8Li with a half-life of 838 ms and 9Li with a half-life of 178.3 ms. All of the remaining radioactive isotopes have half-lives that are shorter than 8.6 ms. The shortest-lived isotope of lithium is 4Li, which decays through proton emission and has a half-life of 7.58043x10−23 s. 7
Li is one of the primordial elements (or, more properly, primordial isotopes) produced in Big Bang nucleosynthesis. A small amount of both 6Li and 7Li are produced in stars, but are thought to be burned as fast as it is produced.[26] Additional small amounts of lithium of both 6Li and 7Li may be generated from solar wind, cosmic rays, and early solar system 7Be and 10Be radioactive decay.[27] 7Li can also be generated in carbon stars.[28] Lithium isotopes fractionate substantially during a wide variety of natural processes,[29] including mineral formation (chemical precipitation), metabolism, and ion exchange. Lithium ions substitute for magnesium and iron in octahedral sites in clay minerals, where 6 Li is preferred to 7Li, resulting in enrichment of the light isotope in processes of hyperfiltration and rock alteration. The exotic 11Li is known to exhibit a nuclear halo.
4
Lithium
Natural occurrence According to theory, the stable isotopes 6Li and 7Li were created in the Big Bang, but the amounts are unclear. Lithium is a fusion fuel in main sequence stars, and there is general agreement that there were larger amounts of lithium in the past than the cosmos contains today. Because of the method by which elements are built up by fusion in stars, there is a general trend in the cosmos Lithium is about as common as chlorine in the Earth's upper that the lighter elements are continental crust, on a per-atom basis. more common. However, lithium (element number 3) is tied with krypton as the 32nd/33rd most abundant element in the cosmos (see Cosmochemical Periodic Table of the Elements in the Solar System), being less common than any element between carbon (element 6) and scandium (element 21). It is not until atomic number 36 (krypton) and beyond that chemical elements are found to be universally less common in the cosmos than lithium. The reasons have to do with the failure of any good mechanisms to synthesize lithium in the fusion reactions between nuclides in supernovae. Due to the absence of any quasi-stable nuclide with five nucleons, nuclei of lithium-5 produced from helium and a proton has no time to fuse with a second proton or neutron to form a six nucleon isotope which might decay to lithium-6, even under extreme conditions of bombardment. Also, the product of helium-helium fusion (berylium-8) is immediately unstable toward disintegration to helium again, and is thus not available for formation of lithium. Some lithium-7 is formed in the pp III branch of the proton-proton chain in main sequence and red giant stars, but it is normally consumed by lithium burning as fast as it is formed. This leaves new formation of the stable isotopes lithium 6 and 7 to rare cosmic ray spallation on carbon or other elements in cosmic dust. Meanwhile, existing Li-6 and Li-7 is destroyed in many nuclear reactions in supernovae and by lithium burning in main sequence stars, resulting in net removal of lithium from the cosmos. In turn the destruction of lithium isotopes is due to their very low energy of binding per nucleon with regard to all other nuclides save deuterium (also destroyed in stars) and helium-3.[30] This low energy of binding encourages breakup of lithium in favor of more tightly-bound nuclides under thermonuclear reaction conditions. Lithium is widely distributed on Earth but does not naturally occur in elemental form due to its high reactivity.[15] Estimates for crustal content range from 20 to 70 ppm by weight.[24] In keeping with its name, lithium forms a minor part of igneous rocks, with the largest concentrations in granites. Granitic pegmatites also provide the greatest abundance of lithium-containing minerals, with spodumene and petalite being the most commercially viable sources.[24] A newer source for lithium is hectorite clay, the only active development of which is through Western Lithium Corp. in the USA. [31]
5
Lithium
6
According to the Handbook of Lithium and Natural Calcium, "Lithium is a comparatively rare element, although it is found in many rocks and some brines, but always in very low concentrations. There are a fairly large number of both lithium mineral and brine deposits but only comparatively a few of them are of actual or potential commercial value. Many are [32] very small, others are too low in grade." At 0.00002 kg lithium per kg of Earth's crust [33] , lithium is the 25th most abundant element. Nickel and lead have the about the same abundance. The largest reserve base of lithium is in the Salar de Uyuni area of Bolivia, which has 5.4 million tons. According to the US Geological Survey, the production and reserves of lithium in metric tons are as follows[34] : Country
2008 Mine Production (est)
Reserves
Reserve Base
Argentina
3200
Not Available
Not Available
Australia
6900
170,000
220,000
Bolivia
None
None
5,400,000
Brazil
180
190,000
910,000
Canada
710
180,000
360,000
Chile
12,000
3,000,000
3,000,000
China
3,500
540,000
1,100,000
Portugal
570
Not Available
Not Available
United States of America
Withheld
38,000
410,000
Zimbabwe
300
23,000
27,000
World Total
27,400
4,100,000
11,000,000
Contrary to the USGS data above, other estimates put Chile's reserve base at 7,520,000 metric tons of lithium, and Argentina's at 6,000,000 metric tons.[35] Seawater contains an estimated 230 billion tons of lithium, though at a low concentration of 0.1 to 0.2 ppm.[36]
Major applications of the metal Because of its specific heat capacity, the highest of all solids, lithium is often used in heat transfer applications. In the latter years of the 20th century lithium became important as an anode material. Used in lithium-ion batteries because of its high electrochemical potential, a typical cell can generate approximately 3 volts, compared with 1.5 volts for lead/acid or zinc cells. Because of its low atomic mass, it also has a high charge- and power-to-weight ratio. Lithium is also used in the pharmaceutical and fine-chemical industry in the manufacture of organolithium reagents, which are used both as strong bases and as reagents for the formation of carbon carbon bonds. Organolithiums are also used in polymer synthesis as catalysts/initiators[37] in anionic polymerisation of unfunctionalised olefins.[38] [39] [40]
Lithium
Medical use Lithium salts were used during the 19th century to treat gout. Lithium salts such as lithium carbonate (Li2CO3), lithium citrate, and lithium orotate are mood stabilizers. They are used in the treatment of bipolar disorder since, unlike most other mood altering drugs, they counteract both mania and depression. Lithium can also be used to augment antidepressants. Because of Lithium's nephrogenic diabetes insipidus effects, it can be used to help treat the syndrome of inappropriate diuretic hormone (SIADH). It was also sometimes prescribed as a preventive treatment for migraine disease and cluster headaches.[41] The active principle in these salts is the lithium ion Li+. Although this ion has a smaller diameter than either Na+ or K+, in a watery environment like the cytoplasmic fluid, Li+ binds to the hydrogen atoms of water, making it effectively larger than either Na+ or K+ ions. How Li+ works in the CNS is still a matter of debate. Li+ elevates brain levels of tryptophan, 5-HT (serotonin), and 5-HIAA (a serotonin metabolite). Serotonin is related to mood stability. Li+ also reduces catecholamine activity in the brain (associated with brain activation and mania), by enhancing reuptake and reducing release. Therapeutically useful amounts of lithium (~ 0.6 to 1.2 mmol/l) are only slightly lower than toxic amounts (>1.5 mmol/l), so the blood levels of lithium must be carefully monitored during treatment to avoid toxicity. Common side effects of lithium treatment include muscle tremors, twitching, ataxia[42] and hypothyroidism. Long term use is linked to hyperparathyroidism[43] , hypercalcemia (bone loss), hypertension, kidney damage, nephrogenic diabetes insipidus (polyuria and polydipsia), seizures[44] and weight gain.[45] Some of the side-effects are a result of the increased elimination of potassium. There appears to be an increased risk of Ebstein (cardiac) Anomaly in infants born to women taking lithium during the first trimester of pregnancy. According to a study in 2009 at Oita University in Japan and published in the British Journal of Psychiatry, communities whose water contained larger amounts of lithium had significantly lower suicide rates[46] [47] [48] [49] but did not address whether lithium in drinking water causes the negative side effects associated with higher doses of the element.[50]
7
Lithium
8
Other uses • Electrical and electronic uses: • Lithium batteries are disposable (primary) batteries with lithium metal or lithium compounds as an anode. Lithium batteries are not to be confused with lithium-ion batteries, which are high energy-density rechargeable batteries. Other rechargeable batteries include the Lithium-ion polymer battery, Lithium iron phosphate battery, and the Nanowire battery. New technologies are constantly being announced. • Lithium niobate is used extensively in telecommunication products such as mobile phones and optical modulators, for such components as resonant crystals. Lithium applications are used in more than 60 percent of mobile phones.[51] • Chemical uses: • Lithium chloride and lithium bromide are extremely hygroscopic and are used as desiccants. • Lithium metal is used in the preparation of organo-lithium compounds. • General engineering:
The red lithium flame leads to Lithium's use in flares and pyrotechnics
• Lithium stearate is a common all-purpose, high-temperature lubricant. • When used as a flux for welding or soldering, lithium promotes the fusing of metals during and eliminates the forming of oxides by absorbing impurities. Its fusing quality is also important as a flux for producing ceramics, enamels and glass. • Alloys of the metal with aluminium, cadmium, copper and manganese are used to make high-performance aircraft parts (see also Lithium-aluminium alloys). • Optics: • Lithium is sometimes used in focal lenses, including spectacles and the glass for the 200-inch (5.08 m) telescope at Mt. Palomar. • The high non-linearity of lithium niobate also makes it useful in non-linear optics applications. • Lithium fluoride, artificially grown as crystal, is clear and transparent and often used in specialist optics for IR, UV and VUV (vacuum UV) applications. It has the lowest refractive index and the farthest transmission range in the deep UV of all common materials. • Rocketry: • Metallic lithium and its complex hydrides, such a Li[AlH4], are used as high energy additives to rocket propellants[3]. • Lithium peroxide, lithium nitrate, lithium chlorate and lithium perchlorate are used as oxidizers in rocket propellants, and also in oxygen candles that supply submarines and space capsules with oxygen.[52] • Nuclear applications: • Lithium deuteride was the fusion fuel of choice in early versions of the hydrogen bomb. When bombarded by neutrons, both 6Li and 7Li produce tritium—this reaction, which was not fully understood when hydrogen bombs were first tested, was responsible for the runaway yield of the Castle Bravo nuclear test. Tritium fuses with deuterium in a
Lithium
9 fusion reaction that is relatively easy to achieve. Although details remain secret, lithium-6 deuteride still apparently plays a role in modern nuclear weapons, as a fusion material.
• Lithium fluoride (highly enriched in the common isotope lithium-7) forms the basic constituent of the preferred fluoride salt mixture (LiF-BeF2) used in liquid-fluoride nuclear reactors. Lithium fluoride is exceptionally chemically stable and LiF/BeF2 mixtures have low melting points and the best neutronic properties of fluoride salt combinations appropriate for reactor use. • In conceptualized nuclear fusion power plants, Lithium will be used to produce tritium in magnetically confined reactors using deuterium and tritium as the fuel. Tritium does not occur naturally and will be produced by surrounding the reacting plasma with a 'blanket' containing lithium where neutrons from the deuterium-tritium reaction in the plasma will react with the lithium to produce more tritium. 6Li + n → 4He + 3H. Various means of doing this will be tested at the ITER reactor being built at Cadarache, France. • Lithium is used as a source for alpha particles, or helium nuclei. When 7Li is bombarded by accelerated protons 8Be is formed, which undergoes spontaneous fission to form two alpha particles. This was the first man-made nuclear reaction, produced by Cockroft and Walton in 1929. • Other uses: • Lithium hydroxide (LiOH) is an important compound of lithium obtained from lithium carbonate (Li2CO3). It is a strong base, and when heated with a fat it produces a lithium soap. Lithium soap has the ability to thicken oils and is used to manufacture lubricating greases. • Lithium hydroxide and lithium peroxide are used in confined areas, such as aboard spacecraft and submarines, for air purification. Lithium hydroxide absorbs carbon dioxide from the air by reacting with it to form lithium carbonate, and is preferred over other alkaline hydroxides for its low weight. Lithium peroxide (Li2O2) in presence of moisture not only absorbs carbon dioxide to form lithium carbonate, but also releases oxygen. For example 2 Li2O2 + 2 CO2 → 2 Li2CO3 + O2. • Lithium compounds are used in red fireworks and flares. • The Mark 50 Torpedo Stored Chemical Energy Propulsion System (SCEPS) uses a small tank of sulfur hexafluoride gas which is sprayed over a block of solid lithium. The reaction generates enormous heat which is used to generate steam from seawater. The steam propels the torpedo in a closed Rankine cycle.[53]
Lithium
10
Production and world supply Since the end of World War II lithium metal production has greatly increased. The metal is separated from other elements in igneous minerals such as those above. Lithium salts are extracted from the water of mineral springs, brine pools and brine deposits. The metal is produced electrolytically from a mixture of fused lithium and potassium chloride. In 1998 it was about US$ 43 per pound ($95 per kg).[54]
Lithium mine, Salar del Hombre Muerto, Argentina. The brine in this salar is rich in lithium, and the mine concentrates the brine by pumping it into solar evaporation ponds. 2009 image from NASA’s EO-1 satellite.
Deposits of lithium are found in South America throughout the Andes mountain chain. Chile is the leading lithium metal producer, followed by Argentina. Both countries recover the lithium from brine pools. In the United States lithium is recovered from brine pools in Nevada.[55] Nearly half the world's known reserves are located in Bolivia, a nation sitting along the central eastern slope of the Andes. In 2009 Bolivia is negotiating with Japanese and French firms to begin extraction.[56] According to the US Geological Survey, Bolivia's Uyuni Desert has 5.4 million tons of lithium, which can be used to make batteries for hybrid and electric vehicles.[56] This is the largest amount of lithium in any country, compared to Chile's 3 million tons and the United States's 760,000 tons.[56] [57] China may emerge as a significant producer of brine-source lithium carbonate around 2010. There is potential production of up to 55,000 tons per year if projects in Qinghai province and Tibet proceed.[58] The total amount of lithium recoverable from global reserves has been estimated at 35 million tonnes, which includes 15 million tons of the known global lithium reserve base.[59] In 1976 a National Research Council Panel estimated lithium resources at 10.6 million tons for the Western World.[60] With the inclusion of Russian and Chinese resources as well as new discoveries in Australia, Serbia, Argentina and the United States, the total had nearly tripled by 2008.[61] [62]
Lithium
11
Precautions
Lithium ingots with a thin layer of black oxide tarnish
Due to its alkaline tarnish, lithium metal is corrosive and requires special handling to avoid skin contact. Breathing lithium dust or lithium compounds (which are often alkaline) initially irritate the nose and throat, while higher exposure can cause a buildup of fluid in the lungs, leading to pulmonary edema. The metal itself is a handling hazard because of the caustic hydroxide produced when it is in contact with moisture. Lithium is safely stored in non-reactive compounds such as naphtha.[63]
Regulation Some jurisdictions limit the sale of lithium batteries, which are the most readily available source of lithium metal for ordinary consumers. Lithium can be used to reduce pseudoephedrine and ephedrine to methamphetamine in the Birch reduction method, which employs solutions of alkali metals dissolved in anhydrous ammonia. Carriage and shipment of some kinds of lithium batteries may be prohibited aboard certain types of transportation (particularly aircraft) because of the ability of most types of lithium batteries to fully discharge very rapidly when short-circuited, leading to overheating and possible explosion in a process called thermal runaway. Most consumer lithium batteries have thermal overload protection built-in to prevent this type of incident, or their design inherently limits short-circuit currents. Internal shorts have been known to develop due to manufacturing defects or damage to batteries that can lead to spontaneous thermal runaway.[64]
See also • Lithium compounds • Lithium-based grease • Dilithium
External links • • • •
International Lithium Alliance [65] USGS: Lithium Statistics and Information WebElements.com – Lithium [67] It's Elemental – Lithium [68]
[66]
• University of Southampton, Mountbatten Centre for International Studies, Nuclear History Working Paper No5. [69]
Lithium
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" Lithium-induced nephrogenic diabetes insipidus (http:/ / www. jabfm. org/ cgi/ content/ abstract/ 12/ 1/ 43)". The Journal of the American Board of Family Practice 12 (1): 43–47. . [45] " Weight Gain and Bipolar Disorder Treatment (http:/ / www. psycheducation. org/ hormones/ Insulin/ weightgain. htm)". PsychEducation.org. November 2007. . [46] " Lithium in drinking water may boost mood (http:/ / www. upi. com/ Science_News/ 2009/ 05/ 01/ Lithium-in-drinking-water-may-boost-mood/ UPI-66841241235675/ )". Science News (United Press International). May 1, 2009 at 11:41 PM. . Retrieved 2009-05-02. [47] Alleyne, Richard (10:01AM BST 01 May 2009). " Natural levels of lithium in drinking water help reduce suicides (http:/ / www. telegraph. co. uk/ health/ healthnews/ 5251365/ Natural-levels-of-lithium-in-drinking-water-help-reduce-suicides. html)". Health: Health News (Telegraph). . Retrieved 2009-05-02. 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Lithium
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[50] " Lithium in water 'curbs suicide' (http:/ / news. bbc. co. uk/ 2/ hi/ health/ 8025454. stm)". Health:Medical Notes. BBC. 09:22 GMT, Friday, 1 May 2009 10:22 UK. . Retrieved 2009-05-02. [51] " You’ve got the power: the evolution of batteries and the future of fuel cells (http:/ / nl. computers. toshiba-europe. com/ Contents/ Toshiba_nl/ NL/ WHITEPAPER/ files/ TISBWhitepapertech. pdf)" (PDF). Toshiba. . Retrieved 2009-05-17. [52] K. Ernst-Christian (2004). "Special Materials in Pyrotechnics: III. Application of Lithium and its Compounds in Energetic Systems". Propellants, Explosives, Pyrotechnics 29 (2): 67–80. doi: 10.1002/prep.200400032 (http:/ / dx. doi. org/ 10. 1002/ prep. 200400032). [53] T.G. Hughes, R.B. Smith, and D.H. Kiely (1983). "Stored Chemical Energy Propulsion System for Underwater Applications". Journal of Energy 7 (2): 128–133. doi: 10.2514/3.62644 (http:/ / dx. doi. org/ 10. 2514/ 3. 62644). [54] Ober, Joyce A. " Lithium (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ lithium/ 450798. pdf)" (pdf). United States Geological Survey. pp. 77-78. . Retrieved 2007-08-19. [55] C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304814. [56] Simon Romero (February 2, 2009). " In Bolivia, a Tight Grip on the Next Big Resource (http:/ / www. nytimes. com/ 2009/ 02/ 03/ world/ americas/ 03lithium. html?ref=world)". New York Times. . [57] " USGS Mineral Commodities Summaries 2009 (http:/ / minerals. usgs. gov/ minerals/ pubs/ mcs/ 2009/ mcs2009. pdf)". USGS. . [58] " The Trouble With Lithium 2 (http:/ / www. meridian-int-res. com/ Projects/ Lithium_Microscope. pdf)" (PDF). Meridian International Research. May 28, 2008. . Retrieved 2008-07-07. [59] " The Trouble with Lithium (http:/ / www. meridian-int-res. com/ Projects/ Lithium_Problem_2. pdf)" (PDF). Meridian International Research. January 2007. . Retrieved 2008-07-07. [60] R.K. Evans (1978). Lithium Reserves and Resources, Energy, Vol 3. Pergamon Press. [61] R.K. Evans (2008). " An Abundance of Lithium (http:/ / www. worldlithium. com/ Abstract. html)". . Retrieved 2009-07-07. [62] R.K. Evans (2008). " An Abundance of Lithium Part 2 (http:/ / www. worldlithium. com/ AN_ABUNDANCE_OF_LITHIUM_-_Part_2. html)". . Retrieved 2009-07-07. [63] A. Keith Furr. (2000). CRC handbook of laboratory safety (http:/ / books. google. de/ books?id=Oo3xAmmMlEwC& pg=PA244). Boca Raton: CRC Press. pp. 244–246. ISBN 9780849325236. . [64] Samuel C. Levy and Per Bro. (1994). Battery hazards and accident prevention (http:/ / books. google. de/ books?id=i7U-0IB8tjMC& pg=PA15). New York: Plenum Press. pp. 15–16. ISBN 9780306447587. . [65] [66] [67] [68] [69]
http:/ / www. lithiumalliance. org/ http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ lithium/ http:/ / www. webelements. com/ lithium/ http:/ / education. jlab. org/ itselemental/ ele003. html http:/ / www. mcis. soton. ac. uk/ Site_Files/ pdf/ nuclear_history/ Working_Paper_No_5. pdf
Article Sources and Contributors
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15
Beryllium
1
Beryllium lithium ← beryllium → boron ↑ Be ↓ Mg
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
4Be Periodic table
Appearance white-gray metallic
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
1462
1608
1791
2023
2327
2742
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1757.1 kJ·mol−1 3rd: 14848.7 kJ·mol−1Atomic radiusCovalent
radius
Beryllium
2
Miscellaneous Crystal structureMagnetic orderingThermal conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusPoisson ratioMohs hardnessVickers hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of beryllium iso 7
Be
9
Be
10
Be
N.A. trace
half-life
DM
53.12 d
100%
9
trace
1.51×106 y
DE (MeV)
DP
ε
0.862
7
γ
0.477
-
β−
0.556
10
Li
Be is stable with 5 neutron B
beryllium, Be, 4 alkaline earth metal2, 2, s9.012182(3) g·mol−1 1s2 2s2 2, 2 (Image) solid 1.85 g·cm−3 1.690 g·cm−3 1560 K,1287 °C,2349 °F 2742 K,2469 °C,4476 °F 7.895 kJ·mol−1 297 kJ·mol−1 (25 °C) 16.443 J·mol−1·K−1 3,[1] 2, 1[2] (amphoteric oxide) 1.57 (Pauling scale) 1st: 899.5 kJ·mol−1112 pm96±3 pm hexagonal diamagnetic (300 K) 200 W·m−1·K−1 (25 °C) 11.3 µm·m−1·K−1 (r.t.) 12870[3] m·s−1 287 GPa 132 GPa 130 GPa 0.032 5.5 1670 MPa 600 MPa 7440-41-7 Beryllium (pronounced /bəˈrɪliəm/) is the chemical element with the symbol Be and atomic number 4. A bivalent element, beryllium is found naturally only combined with other elements in minerals. Notable gemstones which contain beryllium include Beryl (aquamarines and emeralds) and Chrysoberyl (Alexandrite and Cat's eye). The free element is a steel-grey, strong, lightweight brittle alkaline earth metal. It is primarily used as a hardening agent in alloys, notably beryllium copper. Structurally, beryllium's very low density (1.85 times that of water), high melting point (1278 °C), high temperature stability, and low coefficient of thermal expansion, make it in many ways an ideal aerospace material, and it has been used in rocket nozzles and is a significant Beryllium, crystalline fragment component of planned space telescopes. Because of its relatively high transparency to X-rays and other ionizing radiation types, beryllium also has a number of uses as filters and windows for radiation and particle physics experiments. Commercial use of beryllium metal presents technical challenges due to the toxicity (especially by inhalation) of beryllium-containing dusts. Beryllium produces a direct corrosive effect to tissue, and can cause a chronic life-threatening allergic disease called berylliosis in susceptible persons. Beryllium is a relatively rare element in both the Earth and the universe, because it is not formed in conventional stellar nucleosynthesis, but rather during the Big Bang, and later from the action of cosmic rays on interstellar dust. The element is not known to be necessary or useful for either plant or animal life.
Beryllium
History Beryllium was discovered by Louis-Nicolas Vauquelin in 1798 as a component of beryl and in emeralds. Friedrich Wöhler[4] and Antoine Bussy independently isolated the metal in 1828 by reacting potassium and beryllium chloride. Beryllium's chemical similarity to aluminum was probably why beryllium was missed in previous searches.[5]
Etymology The name beryllium comes from the Greek βήρυλλος, bērullos, beryl, from Prakrit veruliya, from Pāli veḷuriya; ] veḷiru or, viḷar, "to become pale," in reference to the pale semiprecious gemstone beryl.[6] For about 160 years, beryllium was also known as glucinium (with the accompanying chemical symbol "Gl"[7] ), the name coming from the Greek word for sweet, due to the sweet taste of its salts.
Characteristics Physical Beryllium has one of the highest melting points of the light metals. It has exceptional elastic rigidity (Young’s modulus 316 GPa). The modulus of elasticity of beryllium is approximately 50% greater than that of steel. The combination of this modulus plus beryllium's relatively low density gives it an unusually fast sound conduction speed at standard conditions (about 12.9 km/s). Other significant properties are the high values for specific heat (1925 J/kg·K) and thermal conductivity (216 W/m·K), which make beryllium the metal with the best heat dissipation characteristics per unit weight. In combination with the relatively low coefficient of linear thermal expansion (11.4 × 10−6 K−1), these characteristics ensure that beryllium demonstrates a unique degree of dimensional stability under conditions of thermal loading.[8] At standard temperature and pressures beryllium resists oxidation when exposed to air (its ability to scratch glass is due to the formation of a thin layer of the hard oxide BeO). It resists corrosion by concentrated nitric acid.[9]
Nuclear Beryllium has a large scattering cross section for high energy neutrons, thus effectively slowing the neutrons to the thermal energy range where the cross section is low (0.008 barn). The predominant beryllium isotope 9Be also undergoes a (n,2n) neutron reaction to 8 Be, i.e. beryllium is a neutron multiplier, releasing more neutrons than it absorbs. Beryllium is highly permeable to X-rays, and neutrons are liberated when it is hit by alpha particles.[8]
3
Beryllium
4
Isotopes Of beryllium's isotopes, only 9Be is stable and the others are relatively unstable or rare. It is thus a mononuclidic element. Cosmogenic 10Be is produced in the atmosphere by cosmic ray spallation of oxygen and nitrogen. Cosmogenic 10Be accumulates at the soil surface, where its relatively long half-life (1.51 million years) permits a long residence time before decaying to 10 Plot showing variations in solar activity, including variation in B. Thus, 10Be and its daughter 10 Be concentration. Note that the beryllium scale is inverted, so products have been used to increases on this scale indictate lower beryllium-10 levels examine soil erosion, soil formation from regolith, the development of lateritic soils, as well as variations in solar activity and the age of ice cores. Solar activity is inversely correlated with Be-10 production, because solar-wind decreases flux of galactic cosmic rays which reach Earth. Beryllium-10 is also formed in nuclear explosions by a reaction of fast neutrons with 13C in the carbon dioxide in air, and is one of the historical indicators of past activity at nuclear test sites.[10] The fact that 7Be and 8Be are unstable has profound cosmological consequences as it means that elements heavier than beryllium could not be produced by nuclear fusion in the Big Bang, since there was insufficient time during the nucleosynthesis phase of the Big Bang expansion to produce carbon by fusion of 4He nuclei and the relatively low concentrations of 8Be available because of its short half-life. Astronomer Fred Hoyle first showed that the energy levels of 8Be and 12C allow carbon production by the triple-alpha process in helium-fueled stars where more synthetic time is available, thus making life possible from the supernova "ash" from these stars. (See also Big Bang nucleosynthesis).[11] 7
Be decays by electron capture, therefore its decay rate is dependent upon its electron configuration - a rare occurrence in nuclear decay.[12] The shortest-lived known isotope of beryllium is 13Be which decays through neutron emission. It has a half-life of 2.7 × 10−21 second. 6Be is also very short-lived with a half-life of 5.0 × 10−21 second.[13] The exotic isotopes
11
Be and
14
Be are known to exhibit a nuclear halo.[14]
Chemical Beryllium has the electronic configuration [He]2s2. In its chemistry Beryllium exhibits the +2 oxidation state and the only evidence of lower valence of beryllium is in the solubility of the metal in BeCl2.[15] The small atomic radius ensures that the Be2+ ion would be highly polarizing leading to significant covalent character in beryllium's bonding.[16] Beryllium is 4 coordinate in complexes e.g. [Be(H2O)4]2+ and tetrahaloberyllates, BeX2−4. This characteristic is used in analytical techniques using EDTA as a ligand which preferentially
Beryllium
5
forms octahedral complexes - thus absorbing other cations such as Al3+ which might interfere, for example in the solvent extraction of a complex formed between Be2+ and acetylacetone.[17] Beryllium metal sits above aluminium in the electrochemical series and would be expected to be a reactive metal, however it is passivated by an oxide layer and does not react with air or water even at red heat.[16] Once ignited however beryllium burns brilliantly forming a mixture of beryllium oxide and beryllium nitride.[16] Beryllium dissolves readily in non-oxidising acids, such as HCl and H2SO4, but not in nitric as this forms the oxide and this behaviour is similar to that of aluminium metal. Beryllium, again similarly to aluminium, dissolves in warm alkali to form the beryllate anion, Be(OH)2−4, and hydrogen gas. The solutions of salts, e.g. beryllium sulfate and beryllium nitrate are acidic because of hydrolysis of the [Be(H2O)4]2+ ion; for example [Be(H2O)4]2+ + H2O
[Be(H2O)3(OH)]+ + H3O+
Compounds Beryllium forms binary compounds with many non-metals. Beryllium hydride is an amorphous white solid believed to be built from corner-sharing {BeH4} tetrahedra.[18] All four anhydrous halides are known. BeF2 has a silica-like structure with corner-shared BeF4 tetrahedra. BeCl2 and BeBr2 have chain structures with edge-shared tetrahedra.[16] They all have linear monomeric gas phase forms.[16] Beryllium oxide, BeO, is a white, high-melting-point solid, which has the wurtzite structure with a thermal conductivity as high as some metals.[16] BeO is amphoteric. Beryllium hydroxide, Be(OH)2 has low solubility in water and is amphoteric.[16] Salts of beryllium can be produced by reacting Be(OH)2 with acid. Beryllium sulfide, selenide and telluride all have the zincblende structure.[15]
Beryllium nitride, Be3N2 is a high-melting-point compound which is readily hydrolysed.[15] Beryllium azide, BeN6 is known and beryllium phosphide, Be3P2 has a similar structure to Be3N2.[15] A number of beryllium borides are known, Be5B, Be4B, Be2B, BeB2, BeB6, BeB12.[15]
Beryllium carbide, Be2C, is a high melting, brick red compound that reacts with water to give methane.[15] No beryllium silicide has been identified.[16] Basic beryllium nitrate and basic beryllium acetate have similar tetrahedral structures with four beryllium atoms coordinated to a central oxide ion.[15]
Occurrence The beryllium content of the earth’s surface rocks is ca. 4 - 6 ppm. Beryllium is a constituent of about 100 out of about 4000 known minerals, the most important of which are bertrandite (Be4Si2O7(OH)2), beryl (Al2Be3Si6O18), chrysoberyl (Al2BeO4), and phenakite (Be2SiO4). Precious forms of beryl are aquamarine, bixbite and emerald.[8]
Production Because of its high affinity for oxygen at elevated temperatures and its ability to reduce water when its oxide film is removed, the extraction of beryllium from its compounds is very difficult. Although electrolysis of a fused mixture of beryllium and sodium fluorides was
Beryllium
6
used to isolate the element in the nineteenth century, the metal's high melting point makes this process more energy intensive than the corresponding production of alkali metals. Early in the twentieth century, the production of beryllium by the thermal decomposition of beryllium iodide was investigated following the success of a similar process for the [19] production of zirconium, but this proved to be uneconomic for volume production. Beryllium metal did not become readily available until 1957. Currently, most is produced by reducing beryllium fluoride with magnesium metal. The price on the US market for vacuum-cast beryllium ingots was 338 US$ per pound ($745/kg) in 2001.[20] BeF2 + Mg → MgF2 + Be
Applications Radiation windows Because of its low atomic number and very
Beryllium target which "converts" a proton beam into a neutron beam
low absorption for X-rays, the oldest and still one of the most important applications of beryllium is in radiation windows for X-ray tubes. Extreme demands are placed on purity and cleanliness of Be to avoid artefacts in the X-ray images. Thin beryllium foils are used as radiation windows for X-ray detectors, and the extremely low absorption minimizes the heating effects caused by high intensity, low energy X-rays typical of synchrotron radiation. Vacuum-tight windows and beam-tubes for radiation experiments on synchrotrons are manufactured exclusively from beryllium. In scientific setups for various X-ray emission studies (e.g., Energy-dispersive X-ray spectroscopy) the sample holder is usually made of beryllium because its emitted X-rays have much lower energies (~100 eV) than X-rays from most studied materials.[8]
Because of its low atomic number beryllium is almost transparent to energetic particles. A square beryllium foil mounted in a steel case to be Therefore it is used to build the beam pipe used as a window between a vacuum chamber and an X-ray microscope. Beryllium, due to its low Z number around the collision region in collider is highly transparent to X-rays. particle physics experiments. Notably all four main detector experiments at the Large Hadron Collider accelerator (ALICE, ATLAS, CMS, LHCb) use a beryllium beam-pipe. Also many high-energy particle physics collision experiments such as the Large Hadron Collider, the Tevatron, the SLAC and others contain beam pipes made of beryllium. Beryllium's low density allows collision products to reach the surrounding detectors without
Beryllium significant interaction, its stiffness allows a powerful vacuum to be produced within the pipe to minimize interaction with gases, its thermal stability allows it to function correctly at temperatures of only a few degrees above absolute zero, and its diamagnetic nature keeps it from interfering with the complex multipole magnet systems used to steer and [21] focus the particle beams.
Mechanical Due to its stiffness, light weight, and dimensional stability over a wide temperature range, beryllium metal is used for lightweight structural components in the defense and aerospace industries in high-speed aircraft, missiles, space vehicles and communication satellites. Several liquid-fueled rockets use nozzles of pure beryllium,[22] [23] Beryllium is used as an alloying agent in the production of beryllium copper, which contains up to 2.5% beryllium. Beryllium-copper alloys are used in many applications because of their combination of high electrical and thermal conductivity, high strength and hardness, nonmagnetic properties, along with good corrosion and fatigue resistance. These applications include the making of spot-welding electrodes, springs, non-sparking tools and electrical contacts. Beryllium was also used in Jason pistols which were used to strip paint from the hulls of ships. In this case, beryllium was alloyed to copper and used as a hardening agent.[24] The excellent elastic rigidity of beryllium has led to its extensive use in precision instrumentation, e.g. in gyroscope inertial guidance systems, and in support structures for optical systems.[8] Beryllium mirrors are a field of particular interest. Large-area mirrors, frequently with a honeycomb support structure, are used, for example, in meteorological satellites where low weight and long-term dimensional stability are critical. Smaller beryllium mirrors are used in optical guidance systems and in fire control systems, e.g. in the German Leopard I and II main battle tanks. In these systems, very rapid movement of the mirror is required which again dictates low mass and high rigidity. Usually the beryllium mirror is coated with hard electroless nickel which can be more easily polished to a finer optical finish than beryllium. In some applications, though, the beryllium blank is polished without any coating. This is particularly applicable to cryogenic operation where thermal expansion mismatch can cause the coating to buckle.[8] The James Webb Space Telescope[25] will have 18 hexagonal beryllium sections for its mirrors. Because JWST will face a temperature of 33 degrees K, the mirror is made of beryllium, capable of handling extreme cold better than glass. Beryllium contracts and deforms less than glass — and remains more uniform — in such temperatures.[26] For the same reason, the optics of the Spitzer Space Telescope are entirely built of beryllium metal.[27] An earlier major application of beryllium was in brakes for military aircraft because of its hardness, high melting point and exceptional heat dissipation. Environmental considerations have led to substitution by other materials.[8] Cross-rolled beryllium sheet is an excellent structural support for printed circuit boards in surface mounted technology. In critical electronic applications, beryllium is both a structural support and heat sink. The application also requires a coefficient of thermal expansion that is well matched to the alumina and polyimide-glass substrates. The beryllium-beryllium oxide composite “E-Materials” have been specially designed for these
7
Beryllium electronic applications and have the additional advantage that the thermal expansion coefficient can be tailored to match diverse substrate materials.[8]
Magnetic • Due to its non-magnetic properties, Beryllium-based tools are often used by military naval EOD-personnel when working on or around sea-mines, as these often have fuses that detonate on direct magnetic contact or when influenced by a magnetic field. • Beryllium-based tools are used for maintenance and construction near MRI scanners. Magnetic tools would be pulled by the scanner's strong magnetic field. Apart from being difficult to remove once magnetic items are stuck in the scanner, the missile-effect can have dangerous consequences. • In the telecommunications industry, tools made of beryllium are used to tune the highly magnetic klystrons used for high power microwave applications.
Nuclear • Beryllium is used in nuclear weapon designs as the outer layer of the pit of the primary stage, surrounding the fissile material. It is a good pusher for implosion, and a very good neutron reflector, as in Beryllium moderated reactors.[28] • Beryllium is sometimes used in neutron sources, in which the beryllium is mixed with an alpha emitter such as 210Po, 226Ra, 239Pu or 241Am.[28] • Beryllium is used in the Joint European Torus fusion research facility and will be used in ITER, to condition the plasma facing components.[29] • Beryllium has also been proposed as a cladding material for nuclear fuel, due to its combination of mechanical, chemical, and nuclear properties.[8]
Acoustics • Beryllium's characteristics (low weight and high rigidity) make it useful as a material for high-frequency drivers. Until recently, most beryllium tweeters used an alloy of beryllium and other metals due to beryllium's high cost and difficulty to form. These challenges, coupled with the high performance of beryllium, caused some manufacturers to falsely claim using pure beryllium.[30] Some high-end audio companies manufacture pure beryllium tweeters or speakers using these tweeters. Because beryllium is many times more expensive than titanium, hard to shape due to its brittleness, and toxicity if mishandled, these tweeters are limited to high-end and public address applications.[31] [32] [33]
Electronic • Beryllium is an effective p-type dopant in III-V compound semiconductors. It is widely used in materials such as GaAs, AlGaAs, InGaAs, and InAlAs grown by molecular beam epitaxy (MBE).[34] • Beryllium oxide is useful for many applications that require the combined properties of an electrical insulator an excellent heat conductor, with high strength and hardness, with a very high melting point. Beryllium oxide is frequently used as an insulator base plate in high-power transistors in RF transmitters for telecommunications. Beryllium oxide is also being studied for use in increasing the thermal conductivity of uranium dioxide nuclear fuel pellets.[35]
8
Beryllium • Beryllium compounds were once used in fluorescent lighting tubes, but this use was discontinued because of berylliosis in the workers manufacturing the tubes.[36]
Toxicity According to the International Agency for Research on Cancer (IARC), beryllium and beryllium compounds are Category 1 carcinogens; they are carcinogenic to both animals and humans.[37] Chronic berylliosis is a pulmonary and systemic granulomatous disease caused by exposure to beryllium. Acute beryllium disease in the form of chemical pneumonitis was first reported in Europe in 1933 and in the United States in 1943. Cases of chronic berylliosis were first described in 1946 Beryllium ore among workers in plants manufacturing fluorescent lamps in Massachusetts. Chronic berylliosis resembles sarcoidosis in many respects, and the differential diagnosis is often difficult. It occasionally killed early workers in nuclear weapons design, such as Herbert Anderson.[38] Although the use of beryllium compounds in fluorescent lighting tubes was discontinued in 1949, potential for exposure to beryllium exists in the nuclear and aerospace industries and in the refining of beryllium metal and melting of beryllium-containing alloys, the manufacturing of electronic devices, and the handling of other beryllium-containing material. Early researchers tasted beryllium and its various compounds for sweetness in order to verify its presence. Modern diagnostic equipment no longer necessitates this highly risky procedure and no attempt should be made to ingest this highly toxic substance. Beryllium and its compounds should be handled with great care and special precautions must be taken when carrying out any activity which could result in the release of beryllium dust (lung cancer is a possible result of prolonged exposure to beryllium laden dust). This substance can be handled safely if certain procedures are followed. No attempt should be made to work with beryllium before familiarization with correct handling procedures. A successful test for beryllium in air and on surfaces has been recently developed and published as a international voluntary consensus standard (ASTM D7202; www.astm.org). The procedure uses dilute ammonium bifluoride for dissolution and fluorescence detection with beryllium bound to sulfonated hydroxybenzoquinoline, allowing detection up to 100 times lower than the recommended limit for beryllium concentration in the workplace. Fluorescence increases with increasing beryllium concentration. The new procedure has been successfully tested on a variety of surfaces and is effective for the dissolution and ultratrace detection of refractory beryllium oxide and silicious beryllium (ASTM D7458).[39]
Inhalation Beryllium is harmful if inhaled and the effects depend on period of exposure. If beryllium concentrations in air are high enough (greater than 100 µg/m³), an acute condition can result, called acute beryllium disease, which resembles pneumonia. Occupational and community air standards are effective in preventing most acute lung damage. Long-term beryllium exposure can increase the risk of developing lung cancer. The more common
9
Beryllium serious health problem from beryllium today is chronic beryllium disease (CBD), discussed below. It continues to occur in industries as diverse as metal recycling, dental laboratories, alloy manufacturing, nuclear weapons production and metal machine shops that work with alloys containing small amounts of beryllium. A 2008 report from the United States National Research Council said that worker exposure to beryllium should be kept "at the lowest feasible level," as the agency's research could not establish any safe level of exposure.[40] Chronic beryllium disease (CBD) Some people (1-15%) are sensitive to beryllium. Sensitization is not an illness, but some of these individuals, if inhaling sufficient quantities of beryllium dust in the micrometer-size range, may have an inflammatory reaction that principally targets the respiratory system and skin. This condition is called chronic beryllium disease (CBD), and can occur within a few months or many years after exposure to higher-than-normal levels of beryllium (greater than 0.2 µg/m³). This disease causes fatigue, weakness, night sweats and can cause difficulty in breathing and a persistent dry cough. It can result in anorexia, weight loss, and may also lead to right-side heart enlargement and heart disease in advanced cases. Some people who are sensitized to beryllium may not have symptoms, and just being sensitized is not a recognized health effect. CBD is treatable, but not curable with traditional drugs and medicine. CBD occurs when the body's immune system recognizes beryllium particles as foreign material and mounts an immune system attack against the particles. Because these particles are typically inhaled into the lungs, the lungs become the major site where the immune system responds. The lung sacs become inflamed and fill with large numbers of white blood cells that accumulate wherever beryllium particles are found. These cells form balls around the beryllium particles called “granulomas.” When enough of these develop, they interfere with the normal function of the organ. Over time, the lungs become stiff and lose their ability to help transfer oxygen from the air into the bloodstream. Patients with CBD develop difficulty inhaling and exhaling sufficient amounts of air, and the amount of oxygen in their bloodstreams falls. Treatment includes supplemental oxygen and immunosuppressants (such as prednisone) to lower the body's overreaction to beryllium. The general population is unlikely to develop acute or chronic beryllium disease because ambient air levels of beryllium are normally very low (<0.03 ng/m3).[41]
Ingestion Swallowing beryllium has not been reported to cause effects in humans because very little beryllium is absorbed from the stomach and intestines. Harmful effects have sometimes been seen in animals ingesting beryllium.[42]
Dermatological effects Beryllium can cause contact dermatitis. Beryllium contact with skin that has been scraped or cut may cause rashes, ulcers, or bumps under the skin called granulomas.[43]
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Beryllium
Effects on children There are no studies on the health effects of children exposed to beryllium, although individual cases of CBD have been reported in children of beryllium workers from the 1940s. It is unknown whether children differ from adults in their susceptibility to beryllium. It is unclear whether beryllium is teratogenic.[44]
Detection in the body Beryllium can be measured in the urine and blood. The amount of beryllium in blood or urine may not indicate time or quantity of exposure. Beryllium levels can also be measured in lung and skin samples. While such measurements may help establish that exposure has occurred, other tests are used to determine if that exposure has resulted in health effects. A blood test, the blood beryllium lymphocyte proliferation test (BeLPT), identifies beryllium sensitization and has predictive value for CBD. The BeLPT has become the standard test for detecting beryllium sensitization and CBD in individuals who are suspected of having CBD and to help distinguish it from similar conditions such as sarcoidosis. It is also the main test used in industry health programs to monitor whether disease is occurring among current and former workers who have been exposed to beryllium on the job. The test can detect disease that is at an early stage, or can detect disease at more advanced stages of illness as well. The BeLPT can also be performed using cells obtained from a person's lung by a procedure called "bronchoscopy".[45]
Industrial release and occupational exposure limits Typical levels of beryllium that industries may release into the air are of the order of 0.01 µg/m³, averaged over a 30-day period, or 2 µg/m³ of workroom air for an 8-hour work shift. Compliance with the current U.S. Occupational Safety and Health Administration (OSHA) permissible exposure limit for beryllium of 2 µg/m³ has been determined to be inadequate to protect workers from developing beryllium sensitization and CBD. The American Conference of Governmental Industrial Hygienists (ACGIH), which is an independent organization of experts in the field of occupational health, has proposed a threshold limit value (TLV) of 0.05 µg/m³ in a 2006 Notice of Intended Change (NIC). This TLV is 40 times lower than the current OSHA permissible exposure limit, reflecting the ACGIH analysis of best available peer-reviewed research data concerning how little airborne beryllium is required to cause sensitization and CBD. Because it can be difficult to control industrial exposures to beryllium, it is advisable to use any methods possible to reduce airborne and surface contamination by beryllium, to minimize the use of beryllium and beryllium-containing alloys whenever possible, and to educate people about the potential hazards if they are likely to encounter beryllium dust or fumes.[46] On 29 January 2009, the Los Alamos National Laboratory announced it was notifying nearly 2,000 current and former employees and visitors that they may have been exposed to beryllium in the lab and may be at risk of disease. Concern over possible exposure to the material was first raised in November 2008, when a box containing beryllium was received at the laboratory's short-term storage facility.[47]
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Beryllium
See also • Category:Beryllium compounds • Sucker Bait, a story by Isaac Asimov in which the health hazard of beryllium dust is an important plot point
References • Burrell, AK. Ehler, DS. McClesky, TM. Minogue, EM. Taylor, TP. Development of a New Fluorescence Method for the Detection of Beryllium on Surfaces. Journal of ASTM International (JAI). 2005. Vol 2: Issue 9. http:/ / www. astm. org/ cgi-bin/ SoftCart. exe/ JOURNALS/ JAI/ PAGES/ JAI13168. htm?E+ mystore • Infante PF, Newman LS. "Commentary: Beryllium exposure and Chronic Beryllium Disease." Lancet 2004; 415-16. • Newman LS. "Beryllium." Chemical & Engineering News, 2003; 36:38. • Kelleher PC, Martyny JW, Mroz MM, Maier LA, Ruttenber JA, Young DA, Newman LS. "Beryllium particulate exposure and disease relations in a beryllium machining plant." J Occup Environ Med 2001; 43:238-249. • Mroz MM, Balkissoon R, Newman LS. "Beryllium." In: Bingham E, Cohrssen B, Powell C (eds.) Patty’s Toxicology, Fifth Edition. New York: John Wiley & Sons 2001, 177-220. • Beryllium and Compounds: TLV Chemical Substances Draft Documentation, Notice of Intended Change ACGIH Publication #7NIC-042
External links • ATSDR Case Studies in Environmental Medicine: Beryllium Toxicity [48] U.S. Department of Health and Human Services • WebElements.com – Beryllium [49] • It's Elemental – Beryllium [50] • National Pollutant Inventory - Beryllium and compounds [51] • MSDS: ESPI Metals [52] • National Institute for Occupational Safety and Health – Beryllium Page [53] • Former Worker Medical Screening Program [54], U.S. Department of Energy • National Supplemental Screening Program (Oak Ridge Associated Universities) [55]
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15
Boron
1
Boron beryllium ← boron → carbon ↑ B ↓ Al
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
5B Periodic table
Appearance black/brown
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties Phase Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
Boron
2 at T/K
2348
2562
2822
3141
3545
4072
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 2427.1 kJ·mol−1 3rd: 3659.7 kJ·mol−1Atomic radiusCovalent radius Miscellaneous Magnetic orderingElectrical resistivityThermal conductivityThermal expansionSpeed of sound (thin rod) Mohs hardnessCAS registry number Most stable isotopes Main article: Isotopes of boron iso
N.A.
half-life
19.9(7)%*
10
11
80.1(7)%*
11
B
DE (MeV)
DP
[1] B is stable with 5 neutron
10
B
DM
[1] B is stable with 6 neutron
*Boron-10 content may be as low as 19.1% and as high as 20.3% in natural samples. Boron-11 is [2] the remainder in such cases.
boron, B, 5 metalloid13, 2, p10.811(7) g·mol−1 [He] 2s2 2p1 2, 3 (Image) solid 2.08 g·cm−3 2349 K,2076 °C,3769 °F 4200 K,3927 °C,7101 °F 50.2 kJ·mol−1 480 kJ·mol−1 (25 °C) 11.087 J·mol−1·K−1 4,[3] 3, 2, 1[4] (mildly acidic oxide) 2.04 (Pauling scale) 1st: 800.6 kJ·mol−190 pm84±3 pm diamagnetic[5] (20 °C) ~106Ω·m (300 K) 27.4 W·m−1·K−1 (25 °C) (ß form) 5–7 [6] µm·m−1·K−1 (20 °C) 16,200 m/s ~9.5 7440-42-8 Boron (pronounced /ˈbɔrɒn/) is the chemical element with atomic number 5 and the chemical symbol B. Boron is a trivalent metalloid element which occurs abundantly in the evaporite ores borax and ulexite. Several allotropes of boron exist: amorphous boron is a brown powder; whereas crystalline boron is black, extremely hard (9.3 on Mohs' scale), and a poor conductor at room temperature. Elemental boron is used as a dopant in the semiconductor industry, while boron compounds play important roles as light structural materials, insecticides and preservatives, and reagents for chemical synthesis. Boron is an essential plant nutrient. Whereas lack of boron results in boron deficiency disorder, high soil concentrations of boron may also be toxic to plants. As an ultratrace element, boron is necessary for the optimal health of rats and presumably other mammals, though its physiological role in animals is not yet fully understood.
Boron
3
Characteristics Allotropes Boron is similar to carbon in its capability to form stable covalently bonded molecular networks. Even nominally disordered (amorphous) boron contains regular boron icosahedra which are, however, bonded randomly to each other without long-range order.[7] [8] Crystalline boron is a very hard, black material with a high melting point of above 2000 °C. It exists in four major polymorphs: α, ß, γ and T. Whereas α, ß and T phases are based on B12 icosahedra, the γ-phase can be described as a rocksalt-type arrangement of the icosahedra and B2 atomic pairs.[9] It can be produced by compressing other boron phases to 12-20 GPa and heating to 1500-1800 °C; it remains stable after releasing the temperature and pressure. The T phase is produced Amorphous boron at similar pressures, but higher temperatures of 1800-2200 °C. As to the α and ß phases, they might both coexist at ambient conditions with the ß phase being more stable.[10] [9] Compressing boron above 160 GPa produces a boron phase with an as yet unknown structure, and this phase is a superconductor at temperatures 6-12 K.[11] Boron phase
α
Symmetry [9]
Atoms/unit cell
[12] [13] [14] [15]
Density (g/cm3)
[16] [17]
Vickers hardness (GPa)
[17] [18]
Bulk modulus (GPa) Bandgap (eV)
[17] [19]
ß
γ
Rhombohedral
Rhombohedral
Orthorhombic
12
~105
28
2.46
2.35
2.52
42
45
50-58
185
224
227
2
1.6
2.1
T Tetragonal
2.36
Chemical Chemically, boron is closer to silicon than to aluminium. Crystalline boron is chemically inert and resistant to attack by boiling hydrofluoric or hydrochloric acid. When finely divided, it is attacked slowly by hot concentrated hydrogen peroxide, hot concentrated nitric acid, hot sulfuric acid or hot mixture of sulfuric and chromic acids.[20] [21] Oxidation of boron depends upon the crystallinity, particle size, purity and temperature. Boron does not react with air at room temperature, but at higher temperatures it burns to form boron trioxide: 4 B + 3 O2 (g) → 2 B2O3 (s) Boron reacts with sulfur to produce boron sulfide: 2 B + 3 S (g) → B2S3 (s) The first synthesis was performed by Jöns Jakob Berzelius in 1824. Another reaction, starting from boron and hydrogen sulfide, was conducted by Friedrich Wöhler and Henri
Boron
4
Etienne Sainte-Claire Deville and published in 1858.[22]
[23]
2 B + 3 H2S → B2S3 (s) + 3 H2 Wöhler and Deville also documented vigorous reactions between boron and the halogens resulting in boron trichloride, boron trifluoride and boron tribromide.[23] For example: 2 B + 3 Br2 → 2 BBr3 Boron can form compounds whose formal oxidation state is not three, such as B(IV) in boron carbide BC,[3] B(II) in B2F4,[24] and B(I) in boron fluoride BF.[4] Boron compounds such as BCl3 behave as electrophiles or Lewis acids in their reactions.[25] Boron is the least electronegative non-metal.[26]
Isotopes Boron has two naturally occurring and stable isotopes, 11B (80.1%) and 10B (19.9%). The mass difference results in a wide range of δ11B values, which are defined as a fractional difference between the 11B and 10B and traditionally expressed in parts per thousand, in natural waters ranging from -16 to +59. There are 13 known isotopes of boron, the shortest-lived isotope is 7B which decays through proton emission and alpha decay. It has a half-life of 3.5×10−22 s. Isotopic fractionation of boron is controlled by the exchange reactions of the boron species B(OH)3 and B(OH)4. Boron isotopes are also fractionated during mineral crystallization, during H2O phase changes in hydrothermal systems, and during hydrothermal alteration of rock. The latter effect results in preferential removal of the 10B(OH)4 ion onto clays. It results in solutions enriched in 11B(OH)3 and therefore may be responsible for the large 11B enrichment in seawater relative to both oceanic crust and continental crust; this difference may act as an isotopic signature. [27] The exotic 17B exhibits a nuclear halo, i.e. its radius is appreciably larger than that predicted by the liquid drop model.[28] Enriched boron (boron-10) The
10
B isotope is good at
capturing thermal neutrons. Natural boron is about 20% 10B and 80%11B. The nuclear industry enriches natural boron to nearly pure 10B. The waste product, or depleted boron, is nearly pure 11B. 11B is a candidate as a fuel for aneutronic fusion and is used in the semiconductor industry. Enriched boron or 10B is used in both radiation shielding and in boron neutron capture therapy. In the latter, a compound containing 10B is attached to a muscle near a
Neutron cross section of boron (top curve is for for 11B)
10
B and bottom curve
Boron tumor. The patient is then treated with a relatively low dose of thermal neutrons. This causes energetic and short range alpha radiation from the boron to bombard the tumor.[29] [30] [31]
In nuclear reactors, 10B is used for reactivity control and in emergency shutdown systems. It can serve either function in the form of borosilicate control rods or as boric acid. In pressurized water reactors, boric acid is added to the reactor coolant when the plant is shut down for refueling. It is then slowly filtered out over many months as fissile material is used up and the fuel becomes less reactive.[32] In future manned interplanetary spacecraft, 10B has a theoretical role as structural material (as boron fibers or BN nanotube material) which would also serve a special role in the radiation shield. One of the difficulties in dealing with cosmic rays, which are mostly high energy protons, is that some secondary radiation from interaction of cosmic rays and spacecraft materials is high energy spallation neutrons. Such neutrons can be moderated by materials high in light elements such as polyethylene, but the moderated neutrons continue to be a radiation hazard unless actively absorbed in the shielding. Among light elements that absorb thermal neutrons, 6Li and 10B appear as potential spacecraft structural materials which serve both for mechanical reinforcement and radiation protection.[33] Depleted boron (boron-11) Cosmic radiation will produce secondary neutrons if it hits spacecraft structures; and neutrons cause fission in 10B if it is present in the spacecraft's semiconductors, producing a gamma ray, an alpha particle, and a lithium ion. The resultant fission products may then dump charge into nearby semiconductor 'chip' structures, causing data loss (bit flipping, or single event upset). In radiation hardened semiconductor designs, one countermeasure is to use depleted boron which is greatly enriched in 11B and contains almost no 10B. 11B is largely immune to radiation damage. Depleted boron is a by-product of the nuclear industry.[32] 11
B is also a candidate as a fuel for aneutronic fusion. When struck by a proton with energy of about 500 keV, it produces three alpha particles and 8.7 MeV of energy. Most other fusion reactions involving hydrogen and helium produce penetrating neutron radiation, which weakens reactor structures and induces long term radioactivity thereby endangering operating personnel. Whereas, the alpha particles from 11B fusion can be turned directly into electric power, and all radiation stops as soon as the reactor is turned off.[34]
5
Boron
6
NMR spectroscopy Both 10B and 11B possess nuclear spin. The nuclear spin of 10B is 3 and that of 11B is 3/2. These isotopes are, therefore, of use in nuclear magnetic resonance spectroscopy; and spectrometers specially adapted to detecting the boron-11 nuclei are available commercially. The 10B and 11B nuclei also cause splitting in the resonances of attached nuclei.[35]
Occurrence
A fragment of ulexite
Boron is a relatively rare element in the Earth's crust, representing only 0.001%. The worldwide commercial borate deposits are estimated as 10 million tonnes.[36] [37] Turkey and the United States are the world's [38] [39] largest producers of boron. Turkey has almost 72% of the world’s boron reserves.[40] Boron does not appear on Earth in elemental form but is found combined in borax, boric acid, colemanite, kernite, ulexite and borates. Boric acid is sometimes found in volcanic spring waters. Ulexite is a borate mineral; it is a fibrous crystal where individual fibers can guide light like optical fibers.[41] Economically important sources of boron are rasorite (kernite) and tincal (borax ore). They are both found in the Mojave Desert of California, but the largest borax deposits are in Central and Western Turkey including the provinces of Eskişehir, Kütahya and Balıkesir [42] [43] [44]
Borax crystals
History and etymology The name boron originates from the Arabic word buraq or the Persian word burah;[45] which are names for the mineral borax.[46] Boron compounds were known thousands of years ago. Borax was known from the deserts of western Tibet, where it received the name of tincal, derived from the Sanskrit. Borax glazes were used in China from AD300, and some tincal even reached the West, where the Arabic alchemist Geber seems to mention it in 700. Marco Polo brought some glazes back to Italy in the 13th century. Agricola, around 1600, reports its use as a flux in metallurgy. In 1777, boric acid was recognized in the hot springs (soffioni) near Florence, Italy, and
Sassolite
Boron became known as sal sedativum, with mainly medical uses. The rare mineral is called sassolite, which is found at Sasso, Italy. This was the main source of European borax from 1827 to 1872, at which date American sources replaced it.[47] [48] Boron was not recognized as an element until it was isolated by Sir Humphry Davy, Joseph Louis Gay-Lussac and Louis Jacques Thénard in 1808 through the reaction of boric acid and potassium. Davy called the element boracium.[49] Jöns Jakob Berzelius identified boron as an element in 1824. The first pure boron was arguably produced by the American chemist W. Weintraub in 1909.[50] [21]
Production Pure elemental boron is difficult to extract. The earliest methods involved reduction of boric oxide with metals such as magnesium or aluminum. However the product is almost always contaminated with metal borides. Pure boron can be prepared by reducing volatile boron halides with hydrogen at high temperatures. Ultrapure boron, for the use in semiconductor industry, is produced by the decomposition of diborane at high temperatures and then further purified with the zone melting or Czochralski processes.[51]
Isotope enrichment Because of its high neutron cross-section, boron-10 is often used to control fission in nuclear reactors as neutron-capturing substance. [52] Several industrial-scale enrichment processes have been developed, however only the fractionated vacuum distillation of the dimethyl ether adduct of boron trifluoride (DME-BF3) and column chromatography of borates are being used. [53]
Market trend Estimated global consumption of boron rose to a record 1.8 million tonnes of B2O3 in 2005, following a period of strong growth in demand from Asia, Europe and North America. Boron mining and refining capacities are considered to be adequate to meet expected levels of growth through the next decade. The form in which boron is consumed has changed in recent years. The use of ores like colemanite has declined following concerns over arsenic content. Consumers have moved towards the use of refined borates and boric acid that have a lower pollutant content. The average cost of crystalline boron is $5/g.[54] Increasing demand for boric acid has led a number of producers to invest in additional capacity. Eti Mine Company of Turkey opened a new boric acid plant with the production capacity of 100,000 tonnes per year at Emet in 2003. Rio Tinto Group increased the capacity of its boron plant from 260,000 tonnes per year in 2003 to 310,000 tonnes per year by May 2005, with plans to grow this to 366,000 tonnes per year in 2006. Chinese boron producers have been unable to meet rapidly growing demand for high quality borates. This has led to imports of disodium tetraborate growing by a hundredfold between 2000 and 2005 and boric acid imports increasing by 28% per year over the same period.[55] [56] The rise in global demand has been driven by high growth rates in fiberglass and borosilicate production. A rapid increase in the manufacture of reinforcement-grade fiberglass in Asia with a consequent increase in demand for borates has offset the development of boron-free reinforcement-grade fiberglass in Europe and the USA. The recent rises in energy prices may lead to greater use of insulation-grade fiberglass, with consequent growth in the boron consumption. Roskill Consulting Group forecasts that
7
Boron
8
world demand for boron will grow by 3.4% per year to reach 21 million tonnes by 2010. The highest growth in demand is expected to be in Asia where demand could rise by an average 5.7% per year.[55] [57]
Applications Glass and ceramics Nearly all boron ore extracted from the Earth is destined for refinement into boric acid and sodium tetraborate. In the United States, 70% of the boron is used for the production of glass and ceramics.[58] Borosilicate glass, which is typically 12%-15% B2O3, 80% SiO2, and 2% Al2O3, has a low coefficient of thermal expansion giving it a good resistance to thermal shock. Duran and Pyrex are two major brand names for this glass.[59] Boron
filaments
are
high-strength,
lightweight
materials that are chiefly used for advanced aerospace structures as a component of composite materials, as well as limited production consumer and sporting goods such as golf clubs and fishing rods.[60] [61] The fibers can be produced by chemical vapor deposition of boron on a tungsten filament.[38] [62] Boron fibers and sub-millimeter sized crystalline boron
Borosilicate glassware. Displayed are two beakers and a test tube.
springs are produced by laser-assisted chemical vapor deposition. Translation of the focused laser beam allows to produce even complex helical structures. Such structures show good mechanical properties (elastic modulus 450 GPa, fracture strain 3.7 %, fracture stress 17 GPa) and can be applied as reinforcement of ceramics or in micromechanical systems.[63]
Semiconductor industry Boron is an important technological dopant for such important semiconductors as silicon, germanium and silicon carbide. Having one less valence electron than the host atom, it donates a hole resulting in p-type conductivity. Traditional method of introducing boron into semiconductors is via its atomic diffusion at high temperatures. This process uses either solid (B2O3), liquid (BBr3) or gaseous boron sources (B2H6 or BF3). However, after 1970s, it was mostly replaced by ion implantation, which relies mostly on BF3 as a boron source.[64] Boron trichloride gas is also an important chemical in semiconductor industry, however not for doping but rather for plasma etching of metals and their oxides.[65]
Boron
9
Engineering materials Boron carbide, a ceramic material which is obtained by decomposing B2O3 with carbon in the electric furnace: 2 B2O3 + 7 C → B4C + 6 CO
Boron carbide is used for inner plates of ballistic vests
It is used in tank armor, bulletproof vests, and numerous other structural applications. Its ability to absorb neutrons without forming long lived radionuclides makes the material attractive as an absorbent for neutron radiation arising in nuclear power plants. Nuclear applications of boron carbide include shielding, control rod and shut down pellets. Within control rods, boron carbide is often powdered, to increase its surface area.[66]
Magnesium diboride is an important superconducting material with the transition temperature of 39 K. MgB2 wires are produced with the powder-in-tube process and applied in superconducting magnets.[67] [68] Boron is a part of neodymium magnet (Nd2Fe14B), which is the strongest type of permanent magnet. It is found in all kinds of domestic and professional electromechanical and electronic devices, such as magnetic resonance imaging (MRI), various motors and actuators, computer HDDs, CD and DVD players, mobile phones, timer switches, speakers, etc.[5]
High-hardness compounds Mechanical properties of BCN solids Material
Diamond
cubic-BC2N
cubic-BC5
cubic-BN
B4C
Vickers hardness (GPa)
115
76
71
62
38
Fracture toughness (MPa m1/2)
5.3
4.5
9.5
6.8
3.5
ReB2 22
Several boron compounds are known for their extreme hardness and toughness, including • Heterodiamond (also called BCN); • Boron nitride. This material is isoelectronic to carbon. Similar to carbon, it has both hexagonal (soft graphite-like h-BN) and cubic (hard, diamond-like c-BN) forms. h-BN is used as a high temperature component and lubricant. c-BN, also known under commercial name borazon,[69] is a superior abrasive. Its hardness is only slightly smaller, but chemical stability is superior to that of diamond. • Rhenium diboride can be produced at ambient pressures, but is rather expensive because of rhenium. The hardness of ReB2 exhibits considerable anisotropy because of its hexagonal layered structure. Its value is comparable to that of tungsten carbide, silicon carbide, titanium diboride or zirconium diboride.[70] • AlMgB14 + TiB2 composites possess high hardness and wear resistance and are used in either bulk form or as coatings for components exposed to high temperatures and wear loads. [71]
Boron Boron carbide and cubic boron nitride powders are widely used as abrasives. Metal borides are used for coating tools through chemical vapor deposition or physical vapor deposition. Implantation of boron ions into metals and alloys, through ion implantation or ion beam deposition, results in a spectacular increase in surface resistance and microhardness. Laser alloying has also been successfully used for the same purpose. These borides are an alternative to diamond coated tools, and their (treated) surfaces have similar properties to those of the bulk boride. [72]
In chemistry Sodium tetraborate pentahydrate (Na2B4O7 • 5H2O) is used in large amounts in making insulating fiberglass and sodium perborate bleach.[73] Sodium tetraborate decahydrate (Na2B4O7 • 10 H2O) can be found in adhesives and in anti-corrosion systems.[74] Sodium borates are used as a flux for soldering silver and gold and with ammonium chloride for welding ferrous metals.[75] They are also fire retarding additives to plastics and rubber articles.[76] Sodium perborate serves as a source of active oxygen in many detergents, laundry detergents, cleaning products, and laundry bleaches. It is also present in some tooth bleaching formulas.[73] Boric acid (also known as orthoboric acid) H3BO3 is used in the production of textile fiberglass and flat panel displays.[77] It also has antiseptic, antifungal, and antiviral properties and for this reasons is applied as a water clarifier in swimming pool water treatment.[78] Boric acid is also traditionally used as an insecticide, notably against ants, fleas, and cockroaches.[79] Triethylborane is a substance which ignites the JP-7 fuel of the Pratt & Whitney J58 turbojet/ramjet engines powering the Lockheed SR-71 Blackbird.[80] It was also used to ignite the F-1 Engines on the Saturn V Rocket utilized by NASA's Apollo and Skylab programs from 1967 until 1973. Triethylborane is suitable for this because of its pyrophoric properties, especially the fact that it burns with very high temperature.[81] Triethylborane is an industrial initiator in radical Navy emergency flare reactions, where it is effective even at low temperatures. It is also injected into vapor deposition reactors as a boron source. Examples are the plasma deposition of boron-containing hard carbon films, silicon nitride-boron nitride films, and for doping of diamond film with boron.[82] Boron compounds show promise in treating arthritis.[83] Because of its distinctive green flame, amorphous boron is used in pyrotechnic flares.[84] It is also used as a melting point depressant in nickel-chromium braze alloys.[85]
10
Boron
11
Biological role There is a boron-containing natural antibiotic, boromycin, isolated from streptomyces.[86] [87] Boron is an essential plant nutrient, required primarily for maintaining the integrity of cell walls. Conversely, high soil concentrations of > 1.0 ppm can cause marginal and tip necrosis in leaves as well as poor overall growth performance. Levels as low as 0.8 ppm can cause these same symptoms to appear in plants particularly sensitive to boron in the soil. Nearly all plants, even those somewhat tolerant of boron in the soil, will show at least some symptoms of boron toxicity when boron content in the soil is greater than 1.8 ppm. When this content exceeds 2.0 ppm, few plants will perform well and some may not survive. When boron levels in plant tissue exceed 200 ppm symptoms of boron toxicity are likely to appear.[88] [89] [90] As an ultratrace element, boron is necessary for the optimal health of rats, although it is necessary in such small amounts that ultrapurified foods and dust filtration of air is necessary to show the effects of boron deficiency, which manifest as poor coat/hair quality. Presumably, boron is necessary to other mammals. No deficiency syndrome in humans has been described. Small amounts of boron occur widely in the diet, and the amounts needed in the diet would, by analogy with rodent studies, is very small. The exact physiological role of boron in the animal kingdom is poorly understood.[91] Boron occurs in all foods produced from plants. Since 1989 its nutritional value has been argued. It is thought that boron plays several biochemical roles in animals, including humans.[92] The U.S. Department of agriculture conducted an experiment in which postmenopausal women took 3 mg of boron a day. The results showed that supplemental boron reduced excretion of calcium by 44%, and activated estrogen and vitamin D. However, whether these effects were conventionally nutritional, or medicinal, could not be determined. The US National Institutes of Health quotes this source: Total daily boron intake in normal human diets ranges from 2.1–4.3 mg boron/kg body weight (bw)/day. [93] [94]
Analytical quantification For determination of boron content in food or materials the colorimetric curcumin method is used. Boron has to be transferred to boric acid or borates and on reaction with curcumin in acidic solution, a red colored boron-chelate complex, rosocyanine, is formed.[95]
Health issues Elemental boron and borates are non-toxic to humans and animals (approximately similar to table salt). The LD50 (dose at which there is 50% mortality) for animals is about 6 g per kg of body weight. Substances with LD50 above 2 g are considered non-toxic. The minimum lethal dose for humans has not been established, but an intake of 4 g/day was reported without incidents, and medical dosages of 20 g of boric acid for neutron capture therapy caused no problems. Fish have survived for 30 min in a saturated boric acid solution and can survive longer in strong borax solutions.[96] Borates are more toxic to insects than to mammals. The boranes and similar gaseous compounds are quite poisonous. As usual, it is not an element that is intrinsically poisonous, but toxicity depends on structure.[47] [48] The boranes (boron hydrogen compounds) are toxic as well as highly flammable and require special care when handling. Sodium borohydride presents a fire hazard due to its reducing
Boron
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nature, and the liberation of hydrogen on contact with acid. Boron halides are corrosive.[97] Congenital endothelial dystrophy type 2, a rare form of corneal dystrophy, is linked to mutations in SLC4A11 gene that encodes a transporter reportedly regulating the intracellular concentration of boron.[98]
See also • • • • • • • • •
Allotropes of boron Category:Boron compounds Boron deficiency Boron oxide Boron nitride Boron neutron capture therapy Boronic acid Hydroboration-oxidation reaction Suzuki coupling
External links • Boron [99] • WebElements.com – Boron [100] • National Pollutant Inventory - Boron and compounds
[101]
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Boron [64] May, Gary S.; Spanos, Costas J. (2006). Fundamentals of semiconductor manufacturing and process control. John Wiley and Sons. p. 51–54. ISBN 0471784060. [65] Sherer, J. Michael (2005). Semiconductor industry: wafer fab exhaust management. CRC Press. p. 39–60. ISBN 1574447203. [66] Weimer, Alan W. (1997). Carbide, Nitride and Boride Materials Synthesis and Processing. Chapman & Hall (London, New York). ISBN 0-412-54060-6. [67] Canfield,, Paul C.; Crabtree, George W. (2003). " Magnesium Diboride: Better Late than Never (http:/ / www. cmp. ameslab. gov/ personnel/ canfield/ pub/ pt0303. pdf)". Physics Today 56 (3): 34–41. doi: 10.1063/1.1570770 (http:/ / dx. doi. org/ 10. 1063/ 1. 1570770). . [68] Braccini, Valeria; Nardelli, Davide; Penco, Roberto; Grasso Giovanni (2007). "Development of ex situ processed MgB2 wires and their applications to magnets". Physica C: Superconductivity 456: 209–217. doi: 10.1016/j.physc.2007.01.030 (http:/ / dx. doi. org/ 10. 1016/ j. physc. 2007. 01. 030). [69] Wentorf, R. H. (1957). "Cubic form of boron nitride". J. Chem Phys. 26: 956. doi: 10.1063/1.1745964 (http:/ / dx. doi. org/ 10. 1063/ 1. 1745964). [70] Qin, Jiaqian (2008). "Is Rhenium Diboride a Superhard Material?". Advanced Materials 20: 4780. doi: 10.1002/adma.200801471 (http:/ / dx. doi. org/ 10. 1002/ adma. 200801471). [71] Schmidt, Jürgen (2007). "Preparation of titanium diboride TiB2 by spark plasma sintering at slow heating rate". Science and Technology of Advanced Materials 8: 376. doi: 10.1016/j.stam.2007.06.009 (http:/ / dx. doi. org/ 10. 1016/ j. stam. 2007. 06. 009). [72] Y. G. Gogotsi and R.A. Andrievski (1999). Materials Science of Carbides, Nitrides and Borides. Springer. pp. 270–270. ISBN 0792357078. [73] C. R. Hammond (2004). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304857. [74] " Borax Decahydrate (http:/ / chemicalland21. com/ industrialchem/ inorganic/ BORAX DECAHYDRATE. htm)". . Retrieved 2009-05-05. [75] A. C. Davies (1992). The Science and Practice of Welding: Welding science and technology. Cambridge University Press. p. 56. ISBN 052143565X. [76] A.R. Horrocks and D. Price (2001). Fire Retardant Materials. Woodhead Publishing Ltd.. p. 55. ISBN 1855734192. [77] F. Ide (2003 url = http:/ / sciencelinks. jp/ j-east/ article/ 200311/ 000020031103A0287941. php). & #32;"Information technology and polymers. Flat panel display". Engineering Materials 51: 84. [78] " Boric acid (http:/ / chemicalland21. com/ industrialchem/ inorganic/ BORIC ACID. htm)". . Retrieved 2009-05-05. [79] Klotz, J. H.; Moss, J. I.; Zhao, R.; Davis, L. R.; Patterson, R. S. (1994). " Oral toxicity of boric acid and other boron compounds to immature cat fleas (Siphonaptera: Pulicidae) (http:/ / grande. nal. usda. gov/ ibids/ index. php?mode2=detail& origin=ibids_references& therow=51171)". J. Econ. Entomol. 87 (6): 1534–1536. . [80] " Lockheed SR-71 Blackbird (http:/ / www. marchfield. org/ sr71a. htm)". March Field Air Museum. . Retrieved 2009-05-05. [81] A. Young (2008). The Saturn V F-1 Engine: Powering Apollo Into History. Springer. p. 86. ISBN 0387096299. [82] Ehrenfried Zschech, Caroline Whelan, Thomas Mikolajick (2005). Materials for information technology: devices, interconnects and packaging. Birkhäuser. p. 44. ISBN 1852339411. [83] Travers, Richard L.; Rennie, George C.; Newnham, Rex E. (1990). "Boron and Arthritis: The Results of a Double-blind Pilot Study". Journal of Nutritional & Environmental Medicine 1 (2): 127–132. doi: 10.3109/13590849009003147 (http:/ / dx. doi. org/ 10. 3109/ 13590849009003147). [84] B. J. Kosanke, B. Sturman, T. Shimizu, I. von Maltitz, R. J. Hancox, M. A. Wilson, N. Kubota, D. R. Dillehay, C. Jennings-White, T. Smith, D. Chapman, M. PodlesakMitwirkende Personen B. J. Kosanke (2004). Pyrotechnic Chemistry. Journal of Pyrotechnics,. pp. 419. ISBN 9781889526157. [85] Wu, Xiaowei; Chandel R. S.; Li, Hang (2001). "Evaluation of transient liquid phase bonding between nickel-based superalloys". Journal of Materials Science 36 (6): 1539–1546. doi: 10.1023/A:1017513200502 (http:/ / dx. doi. org/ 10. 1023/ A:1017513200502). [86] Hütter, R.; Keller-Schien, W.; Knüsel, F.; Prelog, V.; Rodgers jr., G. C.; Suter, P.; Vogel, G.; Voser,W.; Zähner H.; (1967). "Stoffwechselprodukte von Mikroorganismen. 57. Mitteilung. Boromycin". Helvetica Chimica Acta 50: 1533–1539. doi: 10.1002/hlca.19670500612 (http:/ / dx. doi. org/ 10. 1002/ hlca. 19670500612). [87] Dunitz, J. D.; Hawley, D. M. Miklo, D.; White, D. N. J.; Berlin, Yu.; Marui, R.; Prelog, V. (1971). "Structure of boromycin". Helvetica Chimica Acta 54: 1709–1713. doi: 10.1002/hlca.19710540624 (http:/ / dx. doi. org/ 10. 1002/ hlca. 19710540624). [88] Mahler, R. L.. " Essential Plant Micronutrients. Boron in Idaho (http:/ / info. ag. uidaho. edu/ Resources/ PDFs/ CIS1085. pdf)". University of Idaho. . Retrieved 2009-05-05.
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Boron [89] " Functions of Boron in Plant Nutrition (http:/ / www. borax. com/ agriculture/ files/ an203. pdf)" (PDF). U.S. Borax Inc.. . [90] Blevins, Dale G.; Lukaszewski, Krystyna M. (1998). "Functions of Boron in Plant Nutrition". Annual Review of Plant Physiology and Plant Molecular Biology 49: 481–500. doi: 10.1146/annurev.arplant.49.1.481 (http:/ / dx. doi. org/ 10. 1146/ annurev. arplant. 49. 1. 481). [91] Nielsen, Forrest H. (1998). "Ultratrace elements in nutrition: Current knowledge and speculation". The Journal of Trace Elements in Experimental Medicine 11 (2–3): 251–274. doi: 10.1002/(SICI)1520-670X(1998)11:2/3<251::AID-JTRA15>3.0.CO;2-Q (http:/ / dx. doi. org/ 10. 1002/ (SICI)1520-670X(1998)11:2/ 3<251::AID-JTRA15>3. 0. CO;2-Q). [92] " Boron (http:/ / www. pdrhealth. com/ drug_info/ nmdrugprofiles/ nutsupdrugs/ bor_0040. shtml)". PDRhealth. . Retrieved 2008-09-18. [93] Zook, E. G.; Lehman, J. (1965). "Total boron". J. Assoc. Off Agric. Chem 48: 850. [94] United States. Environmental Protection Agency. Office of Water, U. S. Environmental Protection Agency Staff (1993). Health advisories for drinking water contaminants: United States Environmental Protection Agency Office of Water health advisories (http:/ / books. google. de/ books?id=trUdm-GXchIC& pg=PA84). CRC Press. p. 84. ISBN 087371931X. . [95] Silverman, L.; Trego K. (1953). "Corrections-Colorimetric Microdetermination of Boron By The Curcumin-Acetone Solution Method". Anal. Chem. 25: 1639. doi: 10.1021/ac60083a061 (http:/ / dx. doi. org/ 10. 1021/ ac60083a061). [96] Donald E. Garrett (1998). Borates (http:/ / books. google. com/ books?id=imMJJP5T5rsC& pg=PA385). Academic Press. p. 385. ISBN 0122760603. . [97] " Environmental Health Criteria 204: Boron (http:/ / www. inchem. org/ documents/ ehc/ ehc/ ehc204. htm)". the IPCS. 1998. . Retrieved 2009-05-05. [98] Vithana, En; Morgan, P; Sundaresan, P; Ebenezer, Nd; Tan, Dt; Mohamed, Md; Anand, S; Khine, Ko; Venkataraman, D; Yong, Vh; Salto-Tellez, M; Venkatraman, A; Guo, K; Hemadevi, B; Srinivasan, M; Prajna, V; Khine, M; Casey, Jr; Inglehearn, Cf; Aung, T (Jul 2006). "Mutations in sodium-borate cotransporter SLC4A11 cause recessive congenital hereditary endothelial dystrophy (CHED2).". Nature genetics 38 (7): 755–7. doi: 10.1038/ng1824 (http:/ / dx. doi. org/ 10. 1038/ ng1824). ISSN 1061-4036 (http:/ / worldcat. org/ issn/ 1061-4036). PMID 16767101. [99] http:/ / www. du. edu/ ~jcalvert/ phys/ boron. htm [100] http:/ / www. webelements. com/ boron/ [101] http:/ / www. npi. gov. au/ database/ substance-info/ profiles/ 15. html
16
Article Sources and Contributors
Article Sources and Contributors Boron Source: http://en.wikipedia.org/w/index.php?oldid=307040971 Contributors: 2D, 8472, Abradude, Ace of Spades IV, Acroterion, Adrian.benko, Ak BigTrouble, Akldawgs, Alan012, Alanmarkrussell, Alansohn, Alex43223, Alexeymorgunov, AlexiusHoratius, Andre Engels, Andres, Andrew Kanaber, Antandrus, Anwar saadat, Aoganov, Applekid, Archfiendweazal, Archimerged, Arr0n456, Asiir, AtheWeatherman, Ausir, Aussie Alchemist, Avidallred, Avnjay, BRG, Bagatelle, Bantman, Bart133, Barticus88, Bassbonerocks, Beetstra, Bender235, Benjah-bmm27, Beyondthislife, Birdman1, Bkell, Blackangel25, BlueDevil, Bobkeyes, Bomac, Boron1111, Brighterorange, Brockert, Bryan Derksen, Bryce, Bunny Angel13, CWii, CYD, Cacycle, Calabraxthis, Calvin 1998, CanadianLinuxUser, Canley, CapitalR, Capricorn42, Captqrunch, CardinalDan, Carnildo, Cflm001, CharlesC, ChemNerd, Chowbok, Chris Dybala, Christian75, Ciphergoth, Closedmouth, Condem, Conversion script, CopperKettle, Coppertwig, Corbon, Cosmium, Cramapple, Cuchullain, Cyrloc, DRTllbrg, DVD R W, DabMachine, Danski14, Darrien, Darth Panda, David Latapie, DeadEyeArrow, Delldot, Delta G, DennyColt, DerHexer, Dhp1080, Die2u2, Dino, Discospinster, Dlae, Doctorfluffy, Drilnoth, Dwmyers, Dycedarg, ESRFBeam, EdJohnston, Edcolins, Edgar181, Edsanville, El C, Elassint, Eldin raigmore, Eltomzo, Emmyparker, Emperorbma, Epbr123, Eric119, Erik Zachte, Escape Artist Swyer, Evand, FJPB, Facka, FaerieInGrey, Faradayplank, Farosdaughter, Femto, FengRail, Frank Warmerdam, Frankenpuppy, Freedomlinux, Frencheigh, Frymaster, Funeral, Funky Monkey, Fvw, GFZLab, Gamer6484, Gary Cziko, Gene Nygaard, Giftlite, Glenn, Gluckman, Gman124, Grendelkhan, Gryphn, Hak-kâ-ngìn, Hanswaarle, HappyM, HappyVR, Harland1, Haza-w, Hda3ku, Hdt83, Helixblue, HenryLi, Herbee, HexaChord, Hibernian, Hockeyplayer101, HowardJWilk, Icairns, Iridescent, Itub, Ixfd64, J.delanoy, JRM, James A. Stewart, Jan eissfeldt, Jaraalbe, Jeff G., Jeronimo, Jetru, Jjeffrey, John, JohnCD, Jose77, Josh Parris, Jossi, Jrockley, Karl-Henner, Karlhahn, Katalaveno, Ke1v234, Keilana, KnowledgeOfSelf, Komojo, KonradG, Krellis, Ksbrown, Ktsquare, Kurykh, Kusunose, Kwamikagami, L'Aquatique, LA2, Lauriemchorse, Lawrence Cohen, Lazulilasher, LeaveSleaves, Lec CRP1, LightAnkh, LilHelpa, LinaMishima, Ling.Nut, Lon of Oakdale, LorenzoB, LtPowers, LuigiManiac, Lxseto93, MER-C, MPerel, MZMcBride, Madmarigold, Majestik Moose, Mani1, Marcsin, Marnanel, Master of Puppets, Materialscientist, Mav, Mayz, Mazca, McDogm, Mdf, Megaboz, Megan1967, Mgimpel, Mikemill, Milkbreath, Minesweeper, Minnesota1, Miss Madeline, MisterSheik, Mongreilf, Movedgood, Mr0t1633, MrFish, Mufka, Mxn, Mygerardromance, NHRHS2010, Naaa127, Nakon, Nandita 115, Nantoz, Navnløs, NbmMudder, Neillawrence, Nergaal, Netkinetic, Neutrality, Nickptar, Night Gyr, Nihiltres, Nik42, Nilmerg, Nk, Nono64, Noplasma, Numbo3, Ohnoitsjamie, Old Moonraker, Oliver Lineham, Olivier, Orca432, Oxymoron83, PDH, PStatic, Pabouk, Parker2334, PatVanHove, Paul from Michigan, Pb30, Pe1er1, Pegasus1138, Pharaoh of the Wizards, Phgao, Philip Trueman, PierreAbbat, Pikna, Pirkid, Pixel ;-), Pixelface, Plantsurfer, PoliteCarbide, Polonium, Poolkris, Pretzelpaws, Pschemp, Pseudomonas, Psyche825, Puchiko, Pyrotec, Quadell, RTC, RaseaC, Rawling, Red Thunder, RedCoat10, RedWolf, Remember, RexNL, Reyk, Reza kalani, Riana, Rich Farmbrough, Richard Arthur Norton (1958- ), Richardb43, Richfife, Richnotts, Rifleman 82, Rjwilmsi, Robert Foley, Roberta F., Romanm, RoyBoy, Rplix, Rtcoles, Rursus, RyanJones, SEWilco, SH84, SMC, Salem 20078, SamWhitey, Sandahl, Saperaud, Sbharris, Sbmehta, ScaldingHotSoup, Schneelocke, Scott Burley, Sengkang, Shaddack, Shanes, Shirulashem, Sho Uemura, Shrew, Sietse Snel, Skarebo, SkyLined, Sl, Slowking Man, Slucas, Smokefoot, Snaxe920, Soliloquial, Solipsist, Solitude, Someone else, Squids and Chips, Ssbb1234, Staffwaterboy, StephanieM, Stephenb, Steve Crossin, Steven69, Stone, StrontiumDogs, Stui, Stwalkerster, Suisui, Sunborn, Supercoop, Syrthiss, THEN WHO WAS PHONE?, Taqi Haider, Tavilis, Tempodivalse, Tetracube, The Toque, TheSeer, Thingg, Thricecube, Thyraxus, Tim Starling, Tiphareth, Tiptoety, Titoxd, Tlusťa, Tohd8BohaithuGh1, Tom harrison, Travis.Thurston, Trevor MacInnis, Trojancowboy, Trovatore, Ungvichian, Until It Sleeps, Urhixidur, VASANTH S.N., Vcelloho, Vikingforties, Viriditas, Voyagerfan5761, Vsmith, Vssun, Walkerma, Walton One, Warut, Watch37264, Whitepaw, Whkoh, Wikiman7, Wikisaver62, Wknight94, Woohookitty, Wyllium, Xiong Chiamiov, Yamamoto Ichiro, Yyy, Zach4636, Zelator, Zfr, Zotel, 948 anonymous edits
Image Sources, Licenses and Contributors file:Unknown.svg Source: http://en.wikipedia.org/w/index.php?title=File:Unknown.svg License: Public Domain Contributors: Mav file:Electron shell 005 Boron.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_005_Boron.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 file:Boron mNACTEC.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Boron_mNACTEC.jpg License: unknown Contributors: User:Xvazquez File:Bor 1.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Bor_1.jpg License: Creative Commons Attribution-Sharealike 2.0 Contributors: Maksim Image:Neutroncrosssectionboron.png Source: http://en.wikipedia.org/w/index.php?title=File:Neutroncrosssectionboron.png License: Public Domain Contributors: Helix84 Image:ulexita br.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Ulexita_br.jpg License: Public Domain Contributors: Andresboni, Ra'ike Image:Borax crystals.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Borax_crystals.jpg License: unknown Contributors: User:Aramgutang File:Sassolite.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Sassolite.jpg License: unknown Contributors: User:Aramgutang File:Schott Duran glassware.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Schott_Duran_glassware.jpg License: Public Domain Contributors: Original uploader was Skatebiker at en.wikipedia Image:Bodyarmor.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Bodyarmor.jpg License: Public Domain Contributors: Jwissick, 3 anonymous edits Image:Navy emergency flare.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Navy_emergency_flare.jpg License: unknown Contributors: User:Kb
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17
Carbon
1
Carbon boron ← carbon → nitrogen ↑ C ↓ Si
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
6C Periodic table
Appearance Clear (diamond), black (graphite)
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Density (near r.t.) Density (near r.t.) Sublimation pointTriple pointHeat of fusionSpecific heat capacity Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 2352.6 kJ·mol−1 3rd: 4620.5 kJ·mol−1Covalent radiusVan der Waals radius Miscellaneous Magnetic orderingThermal conductivityThermal expansionSpeed of sound
Carbon
2
(thin rod) Young's modulusShear modulusBulk modulusPoisson ratioMohs hardnessCAS registry number Most stable isotopes Main article: Isotopes of carbon iso
N.A.
half-life
12
98.9%
12
13
1.1%
13
14
trace
5730 y
C C C
DM
DE (MeV)
DP 15
C is stable with 6 neutron C is stable with 7 neutron beta-
0.156
14
N
carbon, C, 6 nonmetal14, 2, p12.0107 g·mol−1 1s2 2s2 2p2 or [He] 2s2 2p2 2,4 (Image) Solid amorphous:[1] 1.8 - 2.1 g·cm−3 graphite: 2.267 g·cm−3 diamond: 3.515 g·cm−3 3915 K,3642 °C,6588 °F 4600 K (4327°C), 10800[2] [3] kPa 117 (graphite) kJ·mol−1 (25 °C) 8.517(graphite), 6.155(diamond) J·mol−1·K−14, 3 [4] , 2, 1 [5] , 0, -1, -2, -3, -4[6] 2.55 (Pauling scale) 1st: 1086.5 kJ·mol−1 77(sp³), 73(sp²), 69(sp) pm 170 pm diamagnetic[7] (300 K) 119-165 (graphite) 900-2300 (diamond) W·m−1·K−1 (25 °C) 0.8 (diamond) [8] µm·m−1·K−1 (20 °C) 18350 (diamond) m/s 1050 (diamond) [8] GPa 478 (diamond) [8] GPa 442 (diamond) [8] GPa 0.1 (diamond) [8] 1-2 (Graphite) 10 (Diamond) 7440-44-0 Carbon (pronounced /ˈkɑrbən/) is the chemical element with symbol C and atomic number 6. As a member of group 14 on the periodic table, it is nonmetallic and tetravalent—making four electrons available to form covalent chemical bonds. There are three naturally occurring isotopes, with 12C and 13C being stable, while 14C is radioactive, decaying with a half-life of about 5730 years.[9] Carbon is one of the few elements known since antiquity.[10] [11] The name "carbon" comes from Latin language carbo, coal, and, in some Romance and Slavic languages, the word carbon can refer both to the element and to coal. There are several allotropes of carbon of which the best known are graphite, diamond, and amorphous carbon.[12] The physical properties of carbon vary widely with the allotropic form. For example, diamond is highly transparent, while graphite is opaque and black. Diamond is among the hardest materials known, while graphite is soft enough to form a streak on paper (hence its name, from the Greek word "to write"). Diamond has a very low electrical conductivity, while graphite is a very good conductor. Under normal conditions, diamond has the highest thermal conductivity of all known materials. All the allotropic forms are solids under normal conditions but graphite is the most thermodynamically stable. All forms of carbon are highly stable, requiring high temperature to react even with oxygen. The most common oxidation state of carbon in inorganic compounds is +4, while +2 is found in carbon monoxide and other transition metal carbonyl complexes. The largest sources of inorganic carbon are limestones, dolomites and carbon dioxide, but significant quantities occur in organic deposits of coal, peat, oil and methane clathrates. Carbon forms more compounds than any other element, with almost ten million pure organic compounds described to date, which in turn are a tiny fraction of such compounds that are theoretically possible under standard conditions.[13] Carbon is one of the least abundant elements in the Earth's crust, but the fourth most abundant element in the universe by mass after hydrogen, helium, and oxygen. It is present in all known lifeforms, and in the human body carbon is the second most abundant element
Carbon
3
by mass (about 18.5%) after oxygen.[14] This abundance, together with the unique diversity of organic compounds and their unusual polymer-forming ability at the temperatures commonly encountered on Earth, make this element the chemical basis of all known life.
Characteristics The different forms or allotropes of carbon (see below) include the hardest naturally occurring substance, diamond, and also one of the softest known substances, graphite. Moreover, it has an affinity for bonding with other small atoms, including other carbon atoms, and is capable of forming multiple stable covalent bonds with such atoms. As a result, carbon is known to form almost ten million different compounds; the large majority of all chemical compounds.[13] Carbon also has the highest Theoretically predicted phase diagram melting and sublimation point of all elements. At of carbon atmospheric pressure it has no melting point as its triple point is at 10.8 ± 0.2 MPa and 4600 ± 300 K,[2] [3] [15] [16] so it sublimates at about 3900 K. . Carbon sublimes in a carbon arc which has a temperature of about 5800 K. Thus, irrespective of its allotropic form, carbon remains solid at higher temperatures than the highest melting point metals such as tungsten or rhenium. Although thermodynamically prone to oxidation, carbon resists oxidation more effectively than elements such as iron and copper that are weaker reducing agents at room temperature. Carbon compounds form the basis of all known so far life on Earth, and the carbon-nitrogen cycle provides some of the energy produced by the Sun and other stars. Although it forms an extraordinary variety of compounds, most forms of carbon are comparatively unreactive under normal conditions. At standard temperature and pressure, it resists all but the strongest oxidizers. It does not react with sulfuric acid, hydrochloric acid, chlorine or any alkalis. At elevated temperatures carbon reacts with oxygen to form carbon oxides, and will reduce such metal oxides as iron oxide to the metal. This exothermic reaction is used in the iron and steel industry to control the carbon content of steel: Fe3O4 + 4 C(s) → 3 Fe(s) + 4 CO(g) with sulfur to form carbon disulfide and with steam in the coal-gas reaction C(s) + H2O(g) → CO(g) + H2(g). Carbon combines with some metals at high temperatures to form metallic carbides, such as the iron carbide cementite in steel, and tungsten carbide, widely used as an abrasive and for making hard tips for cutting tools. As to 2009, graphene appears the strongest material ever tested.[17] However, the process of separating it from graphite will require some technological development before it is economical enough to be used in industrial processes.[18] The system of carbon allotropes spans a range of extremes: Synthetic nanocrystalline diamond is the hardest materials known.
Graphite is one of the softest materials known.
Diamond is the ultimate abrasive.
Graphite is a very good lubricant.
Carbon
4
Diamond is an excellent electrical insulator.
Graphite is a conductor of electricity.
Diamond is the best known naturally occurring thermal conductor
Some forms of graphite are used for thermal insulation (i.e. firebreaks and heat shields)
Diamond is highly transparent.
Graphite is opaque.
Diamond crystallizes in the cubic system.
Graphite crystallizes in the hexagonal system.
Amorphous carbon is completely isotropic.
Carbon nanotubes are among the most anisotropic materials ever produced.
Allotropes Atomic carbon is a very short-lived species and therefore, carbon is stabilized in various multi-atomic structures with different molecular configurations called allotropes. The three relatively well-known allotropes of carbon are amorphous carbon, graphite, and diamond. Once considered exotic, fullerenes are nowadays commonly synthesized and used in research; they include buckyballs,[19] [20] carbon nanotubes,[21] carbon nanobuds[22] and nanofibers.[23] [24] Several other exotic allotropes have also been discovered, such as lonsdaleite,[25] glassy carbon,[26] carbon nanofoam[27] and linear acetylenic carbon.[28] • The amorphous form, is an assortment of carbon atoms in a non-crystalline, irregular, glassy state, which is essentially graphite but not held in a crystalline macrostructure. It is present as a powder, and is the main constituent of substances such as charcoal, lampblack (soot) and activated carbon. • At normal pressures carbon takes the form of graphite, in which each atom is bonded trigonally to three others in a plane composed of fused hexagonal rings, just like those in aromatic hydrocarbons. The resulting network is 2-dimensional, and the resulting flat sheets are stacked and loosely bonded through weak Van der Waals forces. This gives graphite its softness and its cleaving properties (the sheets slip easily past one another). Because of the delocalization of one of the outer electrons of each atom to form a π-cloud, graphite conducts electricity, but only in the plane of each covalently bonded sheet. This results in a lower bulk electrical conductivity for carbon than for most metals. The delocalization also accounts for the energetic stability of graphite over diamond at room temperature.
Carbon
• At very high pressures carbon forms the more compact allotrope diamond, having nearly twice the density of graphite. Here, each atom is bonded tetrahedrally to four others, thus making a 3-dimensional network of puckered six-membered rings of atoms. Diamond has the same cubic structure as silicon and germanium and, thanks to the strength of the carbon-carbon bonds is the hardest naturally occurring substance in terms of resistance to scratching. Contrary to the popular belief that "diamonds are forever", they are in fact thermodynamically unstable Some allotropes of carbon: a) diamond; b) graphite; c) lonsdaleite; d–f) under normal conditions fullerenes (C60, C540, C70); g) amorphous carbon; h) carbon nanotube. and transform into graphite.[12] But due to a high activation energy barrier, the transition into graphite is so extremely slow at room temperature as to be unnoticeable. • Under some conditions, carbon crystallizes as lonsdaleite. This form has a hexagonal crystal lattice where all atoms are covalently bonded. Therefore, all properties of lonsdaleite are close to those of diamond. [25] • Fullerenes have a graphite-like structure, but instead of purely hexagonal packing, they also contain pentagons (or even heptagons) of carbon atoms, which bend the sheet into spheres, ellipses or cylinders. The properties of fullerenes (split into buckyballs, buckytubes and nanobuds) have not yet been fully analyzed and represents an intense area of research in nanomaterials. The names "fullerene" and "buckyball" are given after Richard Buckminster Fuller, popularizer of geodesic domes, which resemble the structure of fullerenes. The buckyballs are fairly large molecules formed completely of carbon bonded trigonally, forming spheroids (the best-known and simplest is the soccerball-shaped structure C60 buckminsterfullerene).[19] Carbon nanotubes are structurally similar to buckyballs, except that each atom is bonded trigonally in a curved sheet that forms a hollow cylinder.[20] [21] Nanobuds were first published in 2007 and are hybrid bucky tube/buckyball materials (buckyballs are covalently bonded to the outer wall of a nanotube) that combine the properties of both in a single structure.[22] • Of the other discovered allotropes, Carbon nanofoam is a ferromagnetic allotrope discovered in 1997. It consists of a low-density cluster-assembly of carbon atoms strung together in a loose three-dimensional web, in which the atoms are bonded trigonally in
5
Carbon
6
six- and seven-membered rings. It is among the lightest known solids, with a density of about 2 kg/m³.[29] Similarly, glassy carbon contains a high proportion of closed porosity.[26] But unlike normal graphite, the graphitic layers are not stacked like pages in a book, but have a more random arrangement. Linear acetylenic carbon[28] has the chemical structure[28] -(C:::C)n- .Carbon in this modification is linear with sp orbital hybridization, and is a polymer with alternating single and triple bonds. This type of carbyne is of considerable interest to nanotechnology as its Young's modulus is forty times that of the hardest known material - diamond.[30]
Occurrence An estimate of the global carbon budget: Biosphere, oceans, atmosphere 0.45 x 1018 kilograms Crust Organic carbon
13.2 x 1018 kg
Carbonates
62.4 x 1018 kg Mantle 1200 x 1018 kg
Carbon is the fourth most abundant chemical element in the universe by mass after hydrogen, helium, and oxygen. Carbon is abundant in the Sun, stars, comets, and in the atmospheres of most planets. Some meteorites contain microscopic diamonds that were formed when the solar system was still a protoplanetary disk. Microscopic diamonds may also be formed by the intense pressure and high temperature at the sites of meteorite impacts.[31] Graphite ore
In combination with oxygen in carbon dioxide, carbon is found in the Earth's atmosphere (in quantities of approximately 810 gigatonnes) and dissolved in all water bodies (approximately 36,000 gigatons). Around 1,900 gigatons are present in the biosphere. Hydrocarbons (such as coal, petroleum, and natural gas) contain carbon as well—coal "reserves" (not "resources") amount to around 900 gigatons, and oil reserves around 150 gigatons. With smaller amounts of calcium, magnesium, and iron, carbon is a major component in very large masses of carbonate rock (limestone, dolomite, marble etc.).
Raw diamond crystal.
Carbon
Coal is a significant commercial source of mineral carbon; anthracite containing 92–98% carbon[32] and the largest source (4,000 Gt, or 80% of coal, gas and oil reserves) of carbon in a form suitable for use as fuel.[33] Graphite is found in large quantities in New York and Texas, the United States, Russia, Mexico, Greenland, and India. Natural diamonds occur in the rock kimberlite, found in "Present day" (1990s) sea surface ancient volcanic "necks," or "pipes". Most diamond dissolved inorganic carbon concentration (from the GLODAP deposits are in Africa, notably in South Africa, Namibia, climatology) Botswana, the Republic of the Congo, and Sierra Leone. There are also deposits in Arkansas, Canada, the Russian Arctic, Brazil and in Northern and Western Australia. Diamonds are now also being recovered from the ocean floor off the Cape of Good Hope. However, though diamonds are found naturally, about 30% of all industrial diamonds used in the U.S. are now made synthetically. Carbon-14 is formed in upper layers of the troposphere and the stratosphere, at altitudes of 9–15 km, by a reaction that is precipitated by cosmic rays. Thermal neutrons are produced that collide with the nuclei of nitrogen-14, forming carbon-14 and a proton.
Isotopes Isotopes of carbon are atomic nuclei that contain six protons plus a number of neutrons (varying from 2 to 16). Carbon has two stable, naturally occurring isotopes.[9] The isotope carbon-12 (12C) forms 98.93% of the carbon on Earth, while carbon-13 (13C) forms the remaining 1.07%.[9] The concentration of 12C is further increased in biological materials because biochemical reactions discriminate against 13C.[34] In 1961 the International Union of Pure and Applied Chemistry (IUPAC) adopted the isotope carbon-12 as the basis for atomic weights.[35] Identification of carbon in NMR experiments is done with the isotope 13 C. Carbon-14 (14C) is a naturally occurring radioisotope which occurs in trace amounts on Earth of up to 1 part per trillion (0.0000000001%), mostly confined to the atmosphere and superficial deposits, particularly of peat and other organic materials.[36] This isotope decays by 0.158 MeV β- emission. Because of its relatively short half-life of 5730 years, 14C is virtually absent in ancient rocks, but is created in the upper atmosphere (lower stratosphere and upper troposphere) by interaction of nitrogen with cosmic rays.[37] The abundance of 14C in the atmosphere and in living organisms is almost constant, but decreases predictably in their bodies after death. This principle is used in radiocarbon dating, invented in 1949, which has been used extensively to determine the age of carbonaceous materials with ages up to about 40,000 years.[38] [39] There are 15 known isotopes of carbon and the shortest-lived of these is 8C which decays through proton emission and alpha decay and has a half-life of 1.98739x10-21 s.[40] The exotic 19C exhibits a nuclear halo, which means its radius is appreciably larger than would be expected if the nucleus was a sphere of constant density.[41]
7
Carbon
8
Formation in stars Formation of the carbon atomic nucleus requires a nearly simultaneous triple collision of alpha particles (helium nuclei) within the core of a giant or supergiant star. This happens in conditions of temperature and helium concentration that the rapid expansion and cooling of the early universe prohibited, and therefore no significant carbon was created during the Big Bang. Instead, the interiors of stars in the horizontal branch transform three helium nuclei into carbon by means of this triple-alpha process. In order to be available for formation of life as we know it, this carbon must then later be scattered into space as dust, in supernova explosions, as part of the material which later forms second, third-generation star systems which have planets accreted from such dust. The Solar System is one such third-generation star system. One of the fusion mechanisms powering stars is the carbon-nitrogen cycle. Rotational transitions of various isotopic forms of carbon monoxide (e.g. 12CO, 13CO, and C18O) are detectable in the submillimeter regime, and are used in the study of newly forming stars in molecular clouds.
Carbon cycle Under
terrestrial
conditions,
conversion of one element to another is very rare. Therefore, the amount of carbon on Earth is effectively constant. Thus, processes that use carbon must obtain it somewhere and dispose of it somewhere else. The paths that carbon follows in the environment make up the carbon cycle. For example, plants draw carbon dioxide out of their environment and use it to build biomass, as in carbon respiration or the Calvin cycle, a process of Diagram of the carbon cycle. The black numbers indicate how carbon fixation. Some of this much carbon is stored in various reservoirs, in billions of tons ("GtC" stands for gigatons of carbon; figures are circa 2004). The biomass is eaten by animals, purple numbers indicate how much carbon moves between whereas some carbon is exhaled reservoirs each year. The sediments, as defined in this diagram, by animals as carbon dioxide. The do not include the ~70 million GtC of carbonate rock and carbon cycle is considerably more kerogen. complicated than this short loop; for example, some carbon dioxide is dissolved in the oceans; dead plant or animal matter may become petroleum or coal, which can burn with the release of carbon, should bacteria not consume it.[42]
Carbon
9
Compounds Organic compounds Carbon has the ability to form very long chains of interconnecting C-C bonds. This property is called catenation. Carbon-carbon bonds are strong, and stable. This property allows carbon to form an almost infinite number of compounds; in fact, there are more known carbon-containing compounds than all the compounds of the other chemical elements combined except those of hydrogen (because almost all organic compounds contain hydrogen too).
Structural formula of methane, the simplest possible organic compound.
Correlation between the carbon cycle and formation of organic compounds. In plants, carbon dioxide formed by carbon fixation can join with water in photosynthesis (green) to form organic compounds, which can be utilized and further converted by both plants and animals.
Carbon
10 The simplest form of an organic molecule is the hydrocarbon—a large family of organic molecules that are composed of hydrogen atoms bonded to a chain of carbon atoms. Chain length, side chains and functional groups all affect the properties of organic molecules. By IUPAC's definition, all the other organic compounds are functionalized compounds of hydrocarbons.
Carbon occurs in all known organic life and is the basis of organic chemistry. When united with hydrogen, it forms various flammable compounds called hydrocarbons which are important to industry as refrigerants, lubricants, solvents, as chemical feedstock for the manufacture of plastics and petrochemicals and as fossil fuels. Carbon is the basis for all plastic materials that are used in common household items.
When combined with oxygen and hydrogen, carbon can form many groups of important biological compounds including sugars, lignans, chitins, alcohols, fats, and aromatic esters, carotenoids and terpenes. With nitrogen it forms alkaloids, and with the addition of sulfur also it forms antibiotics, amino acids, and rubber products. With the addition of phosphorus to these other elements, it forms DNA and RNA, the chemical-code carriers of life, and adenosine triphosphate (ATP), the most important energy-transfer molecule in all living cells.
Inorganic compounds Commonly carbon-containing compounds which are associated with minerals or which do not contain hydrogen or fluorine, are treated separately from classical organic compounds; however the definition is not rigid (see reference articles above). Among these are the simple oxides of carbon. The most prominent oxide is carbon dioxide (CO2). This was once the principal constituent of the paleoatmosphere, but is a minor component of the Earth's atmosphere today.[43] Dissolved in water, it forms carbonic acid (H2CO3), but as most compounds with multiple single-bonded oxygens on a single carbon it is unstable.[44] Through this intermediate, though, resonance-stabilized carbonate ions are produced. Some important minerals are carbonates, notably calcite. Carbon disulfide (CS2) is similar. The other common oxide is carbon monoxide (CO). It is formed by incomplete combustion, and is a colorless, odorless gas. The molecules each contain a triple bond and are fairly polar, resulting in a tendency to bind permanently to hemoglobin molecules, displacing oxygen, which has a lower binding affinity.[45] [46] Cyanide (CN–), has a similar structure, but behaves much like a halide ion (pseudohalogen). For example it can form the nitride cyanogen molecule ((CN)2), similar to diatomic halides. Other uncommon oxides are carbon suboxide (C3O2),[47] the unstable dicarbon monoxide (C2O),[48] [49] carbon trioxide (CO3), [50] [51] cyclopentanepentone (C5O5) [52] , cyclohexanehexone (C6O6) [52] , and mellitic anhydride (C12O9). With reactive metals, such as tungsten, carbon forms either carbides (C4–), or acetylides (C2−2) to form alloys with high melting points. These anions are also associated with methane and acetylene, both very weak acids. With an electronegativity of 2.5,[53] carbon prefers to form covalent bonds. A few carbides are covalent lattices, like carborundum (SiC), which resembles diamond.
Carbon
11
Organometallic compounds Organometallic compounds by definition contain at least one carbon-metal bond. A wide range of such compounds exist; major classes include simple alkyl-metal compounds (e.g. tetraethyl lead), η2-alkene compounds (e.g. Zeise's salt, and η3-allyl compounds (e.g. allylpalladium chloride dimer; metallocenes containing cyclopentadienyl ligands (e.g. ferrocene); and transition metal carbene complexes. Many metal carbonyls exist (e.g. tetracarbonylnickel); some workers consider the carbon monoxide ligand to be purely inorganic, and not organometallic. While carbon is understood to exclusively form four bonds, an interesting compound containing an octahedral hexacoordinated carbon atom has been reported. The cation of the compound is [(Ph3PAu)6C]2+. This phenomenon has been attributed to the aurophilicity of the gold ligands.[54]
History and etymology The English name carbon comes from the Latin carbo for coal and charcoal,[55] and hence comes from the French charbon, meaning charcoal. In German, Dutch and Danish, the names for carbon are Kohlenstoff, koolstof and kulstof respectively, all literally meaning coal-substance. Carbon was discovered in prehistory and was known in the forms of soot and charcoal to the earliest human civilizations. Diamonds were known probably as early as 2500 BCE in China, while carbon in the form of charcoal was made around Roman times by the same chemistry as it is today, by heating wood in a pyramid covered with clay to exclude air.[56] [57]
Carl Wilhelm Scheele
Antoine Lavoisier in his youth
In 1722, René A. F. de Réaumur demonstrated that iron was transformed into steel through the absorption of some substance, now known to be carbon.[58] In 1772, Antoine Lavoisier showed that diamonds are a form of carbon, when he burned samples of carbon and diamond then showed that neither produced any water and that both released the same amount of carbon dioxide per gram. Carl Wilhelm Scheele showed that graphite, which had been thought of as a form of lead, was instead a type of carbon.[59] In 1786, the French scientists Claude Louis Berthollet, Gaspard Monge and C. A. Vandermonde then showed that this substance was carbon.[60] In their publication they proposed the name carbone (Latin carbonum) for this element. Antoine Lavoisier listed carbon as an element in his 1789 textbook.[61]
A new allotrope of carbon, fullerene, that was discovered in 1985[62] includes nanostructured forms such as buckyballs and nanotubes.[19] Their discoverers (Curl, Kroto, and Smalley) received the Nobel Prize in Chemistry in 1996.[63] The resulting renewed interest in new forms lead to the discovery of further exotic allotropes, including glassy carbon, and the realization that "amorphous carbon" is not strictly amorphous.[26]
Carbon
Production Graphite Commercially viable natural deposits of graphite occur in many parts of the world, but the most important sources economically are in China, India, Brazil, and North Korea.[64] Graphite deposits are of metamorphic origin, found in association with quartz, mica and feldspars in schists, gneisses and metamorphosed sandstones and limestone as lenses or veins, sometimes of a meter or more in thickness. Deposits of graphite in Borrowdale, Cumberland, England were at first of sufficient size and purity that, until the 1800s, pencils were made simply by sawing blocks of natural graphite into strips before encasing the strips in wood. Today, smaller deposits of graphite are obtained by crushing the parent rock and floating the lighter graphite out on water. According to the USGS, world production of natural graphite in 2006 was 1.03 million tons and in 2005 was 1.04 million tons (revised), of which the following major exporters produced: China produced 720,000 tons in both 2006 and 2005, Brazil 75,600 tons in 2006 and 75,515 tons in 2005 (revised), Canada 28,000 tons in both years, and Mexico (amorphous) 12,500 tons in 2006 and 12,357 tons in 2005 (revised). In addition, there are two specialist producers: Sri Lanka produced 3,200 tons in 2006 and 3,000 tons in 2005 of lump or vein graphite, and Madagascar produced 15,000 tons in both years, a large portion of it "crucible grade" or very large flake graphite. Some other producers produce very small amounts of "crucible grade". According to the USGS, U.S. (synthetic) graphite electrode production in 2006 was 132,000 tons valued at $495 million and in 2005 was 146,000 tons valued at $391 million, and high-modulus graphite (carbon) fiber production in 2006 was 8,160 tons valued at $172 million and in 2005 was 7,020 tons valued at $134 million.
Diamond The diamond supply chain is controlled by a limited number of powerful businesses, and is also highly concentrated in a small number of locations around the world (see figure). Only a very small fraction of the diamond ore consists of actual diamonds. The ore is crushed, during which Diamond output in 2005 care has to be taken in order to prevent larger diamonds from being destroyed in this process and subsequently the particles are sorted by density. Today, diamonds are located in the diamond-rich density fraction with the help of X-ray fluorescence, after which the final sorting steps are done by hand. Before the use of X-rays became commonplace, the separation was done with grease belts; diamonds have a stronger tendency to stick to grease than the other minerals in the ore.[65] Historically diamonds were known to be found only in alluvial deposits in southern India.[66] India led the world in diamond production from the time of their discovery in approximately the 9th century BCE[67] to the mid-18th century AD, but the commercial potential of these sources had been exhausted by the late 18th century and at that time India was eclipsed by Brazil where the first non-Indian diamonds were found in 1725.[68]
12
Carbon
13
Diamond production of primary deposits (kimberlites and lamproites) only started in the 1870s after the discovery of the Diamond fields in South Africa. Production has increased over time and now an accumulated total of 4.5 billion carats have been mined since that date.[69] Interestingly 20% of that amount has been mined in the last 5 years alone and during the last ten years 9 new mines have started production while 4 more are waiting to be opened soon. Most of these mines are located in Canada, Zimbabwe, Angola, and one in Russia.[69] In the United States, diamonds have been found in Arkansas, Colorado, and Montana.[70] [71] In 2004, a startling discovery of a microscopic diamond in the United States[72] led to the January 2008 bulk-sampling of kimberlite pipes in a remote part of Montana.[73] Today, most commercially viable diamond deposits are in Russia, Botswana, Australia and the Democratic Republic of Congo.[74] In 2005, Russia produced almost one-fifth of the global diamond output, reports the British Geological Survey. Australia boasts the richest diamantiferous pipe with production reaching peak levels of 42 metric tons (41 LT; 46 ST) per year in the 1990s.[70] There are also commercial deposits being actively mined in the Northwest Territories of Canada, Siberia (mostly in Yakutia territory, for example Mir pipe and Udachnaya pipe), Brazil, and in Northern and Western Australia. Diamond prospectors continue to search the globe for diamond-bearing kimberlite and lamproite pipes.
Applications Carbon is essential to all known living systems, and without it life as we know it could not exist (see alternative biochemistry). The major economic use of carbon other than food and wood is in the form of hydrocarbons, most notably the fossil fuel methane gas and crude oil (petroleum). Crude oil is used by the petrochemical industry to produce, amongst others, gasoline and kerosene, through a distillation process, in refineries. Cellulose is a natural, carbon-containing polymer produced by plants in the form of cotton, linen, and hemp. Cellulose is mainly used for maintaining structure in plants. Commercially valuable carbon polymers of animal origin include wool, cashmere and silk. Plastics are made from synthetic carbon polymers, often with oxygen and nitrogen atoms included at regular intervals in the main polymer chain. The raw materials for many of these synthetic substances come from crude oil.
Pencil lead for mechanical pencils are made of graphite.
The uses of carbon and its compounds are extremely varied. It can form alloys with iron, of which the most common is carbon steel. Graphite is combined Sticks of vine and compressed charcoal.
Carbon
14 with clays to form the 'lead' used in pencils used for writing and drawing. It is also used as a lubricant and a pigment, as a molding material in glass manufacture, in electrodes for dry batteries and in electroplating and electroforming, in brushes for electric motors and as a neutron moderator in nuclear reactors.
A cloth of woven carbon filaments
Silicon carbide single crystal
Charcoal is used as a drawing material in artwork, for grilling, and in many other uses including iron smelting. Wood, coal and oil are used as fuel for production of energy and space heating. Gem quality diamond is used in jewelry, and Industrial diamonds are used in drilling, cutting and polishing tools for machining metals and stone. Plastics are made from fossil hydrocarbons, and carbon fiber, made by pyrolysis of synthetic polyester fibers is used to reinforce plastics to form advanced, lightweight composite materials. Carbon fiber is made by pyrolysis of extruded and stretched filaments of polyacrylonitrile (PAN) and other organic substances. The crystallographic structure and mechanical properties of the fiber depend on the type of starting material, and on the subsequent processing. Carbon fibers made from PAN have structure resembling narrow filaments of graphite, but thermal processing may re-order the structure into a continuous rolled sheet. The result is fibers with higher specific tensile strength than steel.[75] Carbon black is used as the black pigment in
The C60 fullerene in crystalline form
Tungsten carbide milling bits
printing ink, artist's oil paint and water colours, carbon paper, automotive finishes, India ink and laser printer toner. Carbon black is also used as a filler in rubber products such as tyres and in plastic compounds. Activated charcoal is used as an absorbent and adsorbent in filter material in applications as diverse as gas masks, water purification and kitchen extractor hoods and in medicine to absorb toxins, poisons, or gases from the digestive system. Carbon is used in chemical reduction at high temperatures. Coke is used to reduce iron ore into iron. Case hardening of steel is achieved by heating finished steel components in carbon powder. Carbides of silicon, tungsten, boron and titanium, are among the hardest known
Carbon materials, and are used as abrasives in cutting and grinding tools. Carbon compounds make up most of the materials used in clothing, such as natural and synthetic textiles and leather, and almost all of the interior surfaces in the built environment other than glass, stone and metal.
Diamonds The diamond industry can be broadly separated into two basically distinct categories: one dealing with gem-grade diamonds and another for industrial-grade diamonds. While a large trade in both types of diamonds exists, the two markets act in dramatically different ways. A large trade in gem-grade diamonds exists. Unlike precious metals such as gold or platinum, gem diamonds do not trade as a commodity: there is a substantial mark-up in the sale of diamonds, and there is not a very active market for resale of diamonds. The market for industrial-grade diamonds operates much differently from its gem-grade counterpart. Industrial diamonds are valued mostly for their hardness and heat conductivity, making many of the gemological characteristics of diamond, including clarity and color, mostly irrelevant. This helps explain why 80% of mined diamonds (equal to about 100 million carats or 20,000 kg annually), unsuitable for use as gemstones and known as bort, are destined for industrial use.[76] In addition to mined diamonds, synthetic diamonds found industrial applications almost immediately after their invention in the 1950s; another 3 billion carats (600 metric tons) of synthetic diamond is produced annually for industrial use.[77] The dominant industrial use of diamond is in cutting, drilling, grinding, and polishing. Most uses of diamonds in these technologies do not require large diamonds; in fact, most diamonds that are gem-quality except for their small size, can find an industrial use. Diamonds are embedded in drill tips or saw blades, or ground into a powder for use in grinding and polishing applications.[78] Specialized applications include use in laboratories as containment for high pressure experiments (see diamond anvil cell), high-performance bearings, and limited use in specialized windows.[79] [80] With the continuing advances being made in the production of synthetic diamonds, future applications are beginning to become feasible. Garnering much excitement is the possible use of diamond as a semiconductor suitable to build microchips from, or the use of diamond as a heat sink in electronics.[81]
15
Carbon
Precautions Pure carbon has extremely low toxicity and can be handled and even ingested safely in the form of graphite or charcoal. It is resistant to dissolution or chemical attack, even in the acidic contents of the digestive tract, for example. Consequently if it gets into body tissues it is likely to remain there indefinitely. Carbon black was probably one of the first pigments to be used for tattooing, and Ötzi the Iceman was found to have carbon tattoos that survived during his life and for 5200 years after his death.[82] However, inhalation of coal dust or soot (carbon black) in large quantities can be dangerous, irritating lung tissues and causing the congestive lung disease coalworker's pneumoconiosis. Similarly, diamond dust used as an abrasive can do harm if ingested or inhaled. Microparticles of carbon are produced in diesel engine exhaust Worker at carbon black plant in Sunray, fumes, and may accumulate in the lungs.[83] In these Texas (photo by John Vachon, 1942) examples, the harmful effects may result from contamination of the carbon particles, with organic chemicals or heavy metals for example, rather than from the carbon itself. Carbon may also burn vigorously and brightly in the presence of air at high temperatures, as in the Windscale fire, which was caused by sudden release of stored Wigner energy in the graphite core. Large accumulations of coal, which have remained inert for hundreds of millions of years in the absence of oxygen, may spontaneously combust when exposed to air, for example in coal mine waste tips. The great variety of carbon compounds include such lethal poisons as tetrodotoxin, the lectin ricin from seeds of the castor oil plant Ricinus communis, cyanide (CN-) and carbon monoxide; and such essentials to life as glucose and protein.
See also • • • • •
Carbon chauvinism Carbon footprint Low-carbon economy Organic chemistry Timeline of carbon nanotubes
16
Carbon
17
External links • • • • • • • • • • • •
Carbon - Periodic Table of Videos [84] Carbon on Britannica [85] WebElements.com – Carbon [86] Chemicool.com – Carbon [87] It's Elemental – Carbon [88] Extensive Carbon page at asu.edu [89] Electrochemical uses of carbon [90] Computational Chemistry Wiki [91] Carbon - Super Stuff. Animation with sound and interactive 3D-models. [92] BBC Radio 4 series "In Our Time", on Carbon, the basis of life, 15 June 2006 Introduction to Carbon Properties geared for High School students. [94] Comprehensive Data on Carbon [95]
[93]
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22
Nitrogen
1
Nitrogen carbon ← nitrogen → oxygen ↑ N ↓ P
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
7N Periodic table
Appearance colorless gas General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensityMelting pointBoiling pointTriple pointCritical pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
37
41
46
53
62
77
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 2856 kJ·mol−1 3rd: 4578.1 kJ·mol−1Covalent radiusVan der Waals radius Miscellaneous Crystal structureMagnetic orderingThermal conductivitySpeed of soundCAS registry number Most stable isotopes Main article: Isotopes of nitrogen iso
N.A.
half-life
13
syn
9.965 min
14
99.634%
14
15
0.366%
15
N N N
DM ε
N is stable with 7 neutron N is stable with 8 neutron
DE (MeV) 2.220
DP 13
C
Nitrogen nitrogen, N, 7 nonmetal15, 2, p14.0067(2) g·mol−1 1s2 2s2 2p3 2, 5 (Image) gas (0 °C, 101.325 kPa) 1.251 g/L 63.153 K,-210.00 °C,-346.00 °F 77.36 K,-195.79 °C,-320.3342 °F 63.1526 K (-210°C), 12.53 kPa 126.19 K, 3.3978 MPa (N2) 0.72 kJ·mol−1 (N2) 5.56 kJ·mol−1 (25 °C) (N2) 29.124 J·mol−1·K−1 5, 4, 3, 2, 1, -1, -2, -3 (strongly acidic oxide) 3.04 (Pauling scale) 1st: 1402.3 kJ·mol−171±1 pm 155 pm hexagonal diamagnetic (300 K) 25.83 × 10−3 W·m−1·K−1 (gas, 27 °C) 353 m/s 7727-37-9 Nitrogen (pronounced /ˈnaɪtrədʒɨn/) is a chemical element that has the symbol N and atomic number 7 and atomic mass 14.00674 u. Elemental nitrogen is a colorless, odorless, tasteless and mostly inert diatomic gas at standard conditions, constituting 78% by volume of Earth's atmosphere. Many industrially important compounds, such as ammonia, nitric acid, organic nitrates (propellants and explosives), and cyanides, contain nitrogen. The extremely strong bond in elemental nitrogen dominates nitrogen chemistry, causing difficulty for both organisms and industry in converting the N2 into useful compounds, and releasing large amounts of energy when these compounds burn or decay back into nitrogen gas. The element nitrogen was discovered by Daniel Rutherford, a Scottish physician, in 1772. Nitrogen occurs in all living organisms. It is a constituent element of amino acids and thus of proteins, and of nucleic acids (DNA and RNA). It resides in the chemical structure of almost all neurotransmitters, and is a defining component of alkaloids, biological molecules produced by many organisms.
History Nitrogen (Latin nitrogenium, where nitrum (from Greek nitron) means "saltpetre" (see nitre), and genes means "forming") is formally considered to have been discovered by Daniel Rutherford in 1772, who called it noxious air or fixed air. That there was a fraction of air that did not support combustion was well known to the late 18th century chemist. Nitrogen was also studied at about the same time by Carl Wilhelm Scheele, Henry Cavendish, and Joseph Priestley, who referred to it as burnt air or phlogisticated air. Nitrogen gas was inert enough that Antoine Lavoisier referred to it as "mephetic air" or azote, from the Greek word άζωτος (azotos) meaning "lifeless". Animals died in it, and it was the principal component of air in which animals had suffocated and flames had burned to extinction. Lavoisier's name for nitrogen is used in many languages (French, Russian, etc.) and still remains in English in the common names of many compounds, such as hydrazine and compounds of the azide ion. Compounds of nitrogen were known in the Middle Ages. The alchemists knew nitric acid as aqua fortis (strong water). The mixture of nitric and hydrochloric acids was known as aqua regia (royal water), celebrated for its ability to dissolve gold (the king of metals). The earliest military, industrial and agricultural applications of nitrogen compounds involved uses of saltpeter (sodium nitrate or potassium nitrate), notably in gunpowder, and much later, as fertilizer.
2
Nitrogen
Properties Nitrogen is a nonmetal, with an electronegativity of 3.04. It has five electrons in its outer shell and is therefore trivalent in most compounds. The triple bond in molecular nitrogen (N2) is the strongest in nature. The resulting difficulty of converting N2 into other compounds, and the ease (and associated high energy release) of converting nitrogen compounds into elemental N2, have dominated the role of nitrogen in both nature and human economic activities. At atmospheric pressure molecular nitrogen condenses (liquifies) at 77 K (−195.8 °C) and freezes at 63 K (−210.0 °C) into the beta hexagonal close-packed crystal allotropic form. Below 35.4 K (−237.6 °C) nitrogen assumes the alpha cubic crystal allotropic form. Liquid nitrogen, a fluid resembling water in appearance, but with 80.8% of the density (the density of liquid nitrogen at its boiling point is 0.808 g/mL), is a common cryogen. Unstable allotropes of nitrogen consisting of more than two nitrogen atoms have been produced in the laboratory, like N3 and N4.[1] Under extremely high pressures (1.1 million atm) and high temperatures (2000 K), as produced using a diamond anvil cell, nitrogen polymerizes into the single-bonded cubic gauche crystal structure. This structure is similar to that diamond, and both have extremely strong covalent bonds. N4 is nicknamed "nitrogen diamond."[2]
Isotopes There are two stable isotopes of nitrogen: 14N and 15N. By far the most common is 14N (99.634%), which is produced in the CNO cycle in stars. Of the ten isotopes produced synthetically, 13N has a half-life of ten minutes and the remaining isotopes have half-lives on the order of seconds or less. Biologically-mediated reactions (e.g., assimilation, nitrification, and denitrification) strongly control nitrogen dynamics in the soil. These reactions typically result in 15N enrichment of the substrate and depletion of the product. 0.73% of the molecular nitrogen in Earth's atmosphere is comprised of the isotopologue 14 15 N N and almost all the rest is 14N2.
Radioisotope 16N is the dominant radionuclide in the coolant of pressurized water reactors during normal operation. It is produced from 16O (in water) via (n,p) reaction. It has a short half-life of about 7.1 s, but during its decay back to 16O produces high-energy gamma radiation (5 to 7 MeV). Because of this, the access to the primary coolant piping must be restricted during reactor power operation[3] . 16N is one of the main means used to immediately detect even small leaks from the primary coolant to the secondary steam cycle.
Electromagnetic spectrum Molecular nitrogen (14N2) is largely transparent to infrared and visible radiation because it is a homonuclear molecule and thus has no dipole moment to couple to electromagnetic radiation at these wavelengths. Significant absorption occurs at extreme ultraviolet wavelengths, beginning around 100 nanometers. This is associated with electronic transitions in the molecule to states in which charge is not distributed evenly between nitrogen atoms. Nitrogen absorption leads to significant absorption of ultraviolet radiation in the Earth's upper atmosphere as well as in the atmospheres of other planetary bodies. For similar reasons, pure molecular nitrogen lasers typically emit light in the ultraviolet range.
3
Nitrogen
4
Nitrogen also makes a contribution to visible air glow from the Earth's upper atmosphere, through electron impact excitation followed by emission. This visible blue air glow (seen in the polar aurora and in the re-entry glow of returning spacecraft) typically results not from molecular nitrogen, but rather from free nitrogen atoms combining with oxygen to form nitric oxide (NO).
Reactions Nitrogen is generally unreactive at standard temperature and pressure. N2 reacts spontaneously with few reagents, being resilient to acids and bases as well as oxidants and most reductants. When nitrogen reacts spontaneously with a reagent, the net transformation is often called nitrogen fixation. Nitrogen reacts with elemental lithium at STP.[4] Lithium burns in an atmosphere of N2 to give lithium nitride: 6 Li + N2 → 2 Li3N Magnesium also burns in nitrogen, forming magnesium nitride.
Structure of [Ru(NH3)5(N2)]2+.
3 Mg + N2 → Mg3N2 N2 forms a variety of adducts with transition metals. The first example of a dinitrogen complex is [Ru(NH3)5(N2)]2+ (see figure at right). Such compounds are now numerous, other examples include IrCl(N2)(PPh3)2, W(N2)2(Ph2CH2CH2PPh2)2, and 5 [(η -C5Me4H)2Zr]2(μ2,η²,η²-N2). These complexes illustrate how N2 might bind to the metal(s) in nitrogenase and the catalyst for the Haber process.[5] A catalytic process to reduce N2 to ammonia with the use of a molybdenum complex in the presence of a proton source was published in 2005.[4] (see nitrogen fixation) The starting point for industrial production of nitrogen compounds is the Haber process, in which nitrogen is fixed by reacting N2 and H2 over an iron(III) oxide (Fe3O4) catalyst at about 500 °C and 200 atmospheres pressure. Biological nitrogen fixation in free-living cyanobacteria and in the root nodules of plants also produces ammonia from molecular nitrogen. The reaction, which is the source of the bulk of nitrogen in the biosphere, is catalyzed by the nitrogenase enzyme complex which contains Fe and Mo atoms, using energy derived from hydrolysis of adenosine triphosphate (ATP) into adenosine diphosphate and inorganic phosphate (−20.5 kJ/mol).
Occurrence Nitrogen is the largest single constituent of the Earth's atmosphere (78.082% by volume of dry air, 75.3% by weight in dry air). It is created by fusion processes in stars, and is estimated to be the 7th most abundant chemical element by mass in the universe.[6] Molecular nitrogen and nitrogen compounds have been detected in interstellar space by astronomers using the Far Ultraviolet Spectroscopic Explorer.[7] Molecular nitrogen is a major constituent of the Saturnian moon Titan's thick atmosphere, and occurs in trace amounts in other planetary atmospheres.[8]
Nitrogen Nitrogen is present in all living organisms, in proteins, nucleic acids and other molecules. It typically makes up around 4% of the dry weight of plant matter, and around 3% of the weight of the human body. It is a large component of animal waste (for example, guano), usually in the form of urea, uric acid, ammonium compounds and derivatives of these nitrogenous products, which are essential nutrients for all plants that are unable to fix atmospheric nitrogen. Nitrogen occurs naturally in a number of minerals, such as saltpetre (potassium nitrate), Chile saltpetre (sodium nitrate) and sal ammoniac (ammonium chloride). Most of these are relatively uncommon, partly because of the minerals' ready solubility in water. See also Nitrate minerals and Ammonium minerals.
Compounds The main neutral hydride of nitrogen is ammonia (NH3), although hydrazine (N2H4) is also commonly used. Ammonia is more basic than water by 6 orders of magnitude. In solution ammonia forms the ammonium ion (NH+4). Liquid ammonia (boiling point 240 K) is amphiprotic (displaying either Brønsted-Lowry acidic or basic character) and forms ammonium and the less common amide ions (NH−2); both amides and nitride (N3−) salts are known, but decompose in water. Singly, doubly, triply and quadruply substituted alkyl compounds of ammonia are called amines (four substitutions, to form commercially and biologically important quaternary amines, results in a positively charged nitrogen, and thus a water-soluble, or at least amphiphilic, compound). Larger chains, rings and structures of nitrogen hydrides are also known, but are generally unstable. N2+2 is another polyatomic cation as in hydrazine. Other classes of nitrogen anions (negatively charged ions) are the poisonous azides (N−3), which are linear and isoelectronic to carbon dioxide, but which bind to important iron-containing enzymes in the body in a manner more resembling cyanide. Another molecule of the same structure is the colorless and relatively inert anesthetic gas Nitrous oxide (dinitrogen monoxide, N2O), also known as laughing gas. This is one of a variety of nitrogen oxides that form a family often abbreviated as NOx. Nitric oxide (nitrogen monoxide, NO), is a natural free radical used in signal transduction in both plants and animals, for example in vasodilation by causing the smooth muscle of blood vessels to relax. The reddish and poisonous nitrogen dioxide NO2 contains an unpaired electron and is an important component of smog. Nitrogen molecules containing unpaired electrons show an understandable tendency to dimerize (thus pairing the electrons), and are generally highly reactive. The corresponding acids are nitrous HNO2 and nitric acid HNO3, with the corresponding salts called nitrites and nitrates. The higher oxides dinitrogen trioxide N2O3, dinitrogen tetroxide N2O4 and dinitrogen pentoxide N2O5, are fairly unstable and explosive, a consequence of the chemical stability of N2. N2O4 is one of the most important oxidizers of rocket fuels, used to oxidize hydrazine in the Titan rocket and in the recent NASA MESSENGER probe to Mercury. N2O4 is an intermediate in the manufacture of nitric acid HNO3, one of the few acids stronger than hydronium and a fairly strong oxidizing agent.
5
Nitrogen
6
Nitrogen is notable for the range of explosively unstable compounds that it can produce. Nitrogen triiodide NI3 is an extremely sensitive contact explosive. Nitrocellulose, produced by nitration of cellulose with nitric acid, is also known as guncotton. Nitroglycerin, made by nitration of glycerin, is the dangerously unstable explosive ingredient of dynamite. The comparatively stable, but more powerful explosive trinitrotoluene (TNT) is the standard explosive against which the power of nuclear explosions are measured. Nitrogen can also be found in organic compounds. Common nitrogen functional groups include: amines, amides, nitro groups, imines, and enamines. The amount of nitrogen in a chemical substance can be determined by the Kjeldahl method.
Applications Nitrogen gas is an industrial gas produced by the fractional distillation of liquid air, or by mechanical means using gaseous air (i.e. pressurized reverse osmosis membrane or Pressure swing adsorption). Commercial nitrogen is often a byproduct of air-processing for industrial concentration of oxygen for steelmaking and other purposes. When supplied compressed in cylinders it is often referred to as OFN (oxygen-free nitrogen).[9] Nitrogen gas has a wide variety of applications, including serving as an inert replacement for air where oxidation is undesirable;
A computer rendering of the nitrogen molecule, N2.
• To preserve the freshness of packaged or bulk foods (by delaying rancidity and other forms of oxidative damage) • • • • • •
In ordinary incandescent light bulbs as an inexpensive alternative to argon.[10] On top of liquid explosives as a safety measure The production of electronic parts such as transistors, diodes, and integrated circuits Dried and pressurized, as a dielectric gas for high voltage equipment The manufacturing of stainless steel Use in military aircraft fuel systems to reduce fire hazard, (see inerting system)
• Filling automotive and aircraft tires[11] due to its inertness and lack of moisture or oxidative qualities, as opposed to air, though this is not necessary for consumer automobiles.[12] [13] Nitrogen molecules are less likely to escape from the inside of a tire compared with the traditional air mixture used. Air consists mostly of nitrogen and oxygen. Nitrogen molecules have a larger effective diameter than oxygen molecules and therefore diffuse through porous substances more slowly.[14] Nitrogen is commonly used during sample preparation procedures for chemical analysis. Specifically, it is used as a means of concentrating and reducing the volume of liquid samples. Directing a pressurized stream of nitrogen gas perpendicular to the surface of the liquid allows the solvent to evaporate while leaving the solute(s) and un-evaporated solvent behind.[15] Nitrogen tanks are also replacing carbon dioxide as the main power source for paintball guns. The downside is that nitrogen must be kept at higher pressure than CO2, making N2 tanks heavier and more expensive.
Nitrogen
Nitrogenated beer A further example of its versatility is its use as a preferred alternative to carbon dioxide to pressurize kegs of some beers, particularly stouts and British ales, due to the smaller bubbles it produces, which make the dispensed beer smoother and headier. A modern application of a pressure sensitive nitrogen capsule known commonly as a "widget" now allows nitrogen charged beers to be packaged in cans and bottles.[16]
Liquid nitrogen Liquid nitrogen is a cryogenic liquid. At atmospheric pressure, it boils at −195.8 °C. When insulated in proper containers such as Dewar flasks, it can be transported without much evaporative loss. Like dry ice, the main use of liquid nitrogen is as a refrigerant. Among other things, it is used in the cryopreservation of blood, reproductive cells (sperm and egg), and other biological samples and materials. It is used in cold traps for certain laboratory equipment and to cool x-ray detectors. It has also been used to cool central processing units and other devices in computers which are overclocked, and which produce more heat than during normal operation.
Applications of nitrogen compounds Molecular nitrogen (N2) in the atmosphere is relatively non-reactive due to its strong bond, and N2 plays an inert role in the human body, being neither produced nor destroyed. In nature, nitrogen is converted into biologically (and industrially) useful compounds by lightning, and by some living organisms, notably certain bacteria (i.e. nitrogen fixing bacteria – see Biological role below). Molecular nitrogen is released into the atmosphere in the process of decay, in dead plant and animal tissues. The ability to combine or fix molecular nitrogen is a key feature of modern industrial chemistry, where nitrogen and natural gas are converted into ammonia via the Haber process. Ammonia, in turn, can be used directly (primarily as a fertilizer, and in the synthesis of nitrated fertilizers), or as a precursor of many other important materials including explosives, largely via the production of nitric acid by the Ostwald process. The organic and inorganic salts of nitric acid have been important historically as convenient stores of chemical energy. They include important compounds such as potassium nitrate (or saltpeter used in gunpowder) and ammonium nitrate, an important fertilizer and explosive (see ANFO). Various other nitrated organic compounds, such as nitroglycerin and trinitrotoluene, and nitrocellulose, are used as explosives and propellants for modern firearms. Nitric acid is used as an oxidizing agent in liquid fueled rockets. Hydrazine and hydrazine derivatives find use as rocket fuels and monopropellants. In most of these compounds, the basic instability and tendency to burn or explode is derived from the fact that nitrogen is present as an oxide, and not as the far more stable nitrogen molecule (N2) which is a product of the compounds' thermal decomposition. When nitrates burn or explode, the formation of the powerful triple bond in the N2 produces most of the energy of the reaction. Nitrogen is a constituent of molecules in every major drug class in pharmacology and medicine. Nitrous oxide (N2O) was discovered early in the 19th century to be a partial anesthetic, though it was not used as a surgical anesthetic until later. Called "laughing gas", it was found capable of inducing a state of social disinhibition resembling
7
Nitrogen drunkenness. Other notable nitrogen-containing drugs are drugs derived from plant alkaloids, such as morphine (there exist many alkaloids known to have pharmacological effects; in some cases they appear natural chemical defenses of plants against predation). Nitrogen containing drugs include all of the major classes of antibiotics, and organic nitrate drugs like nitroglycerin and nitroprusside which regulate blood pressure and heart action by mimicking the action of nitric oxide.
Biological role Nitrogen is an essential building block of amino and nucleic acids, essential to life on Earth. Elemental nitrogen in the atmosphere cannot be used directly by either plants or animals, and must converted to a reduced (or 'fixed') state in order to be useful for higher plants and animals. Precipitation often contains substantial quantities of ammonium and nitrate, thought to result from nitrogen fixation by lightning and other atmospheric electric phenomena.[17] This was first proposed by Liebig in 1827 and later confirmed.[17] However, because ammonium is preferentially retained by the forest canopy relative to atmospheric nitrate, most fixed nitrogen that reaches the soil surface under trees as nitrate. Soil nitrate is preferentially assimilated by these tree roots relative to soil ammonium. Specific bacteria (e.g. Rhizobium trifolium) possess nitrogenase enzymes which can fix atmospheric nitrogen (see nitrogen fixation) into a form (ammonium ion) that is chemically useful to higher organisms. This process requires a large amount of energy and anoxic conditions. Such bacteria may live freely in soil (e.g. Azotobacter) but normally exist in a symbiotic relationship in the root nodules of leguminous plants (e.g. clover, Trifolium, or soybean plant, Glycine max). Nitrogen-fixing bacteria are also symbiotic with a number of unrelated plant species such as alders (Alnus) spp., lichens (Casuarina), Myrica, liverworts, and Gunnera. As part of the symbiotic relationship, the plant converts the 'fixed' ammonium ion to nitrogen oxides and amino acids to form proteins and other molecules, (e.g. alkaloids). In return for the 'fixed' nitrogen, the plant secretes sugars to the symbiotic bacteria. Some plants are able to assimilate nitrogen directly in the form of nitrates which may be present in soil from natural mineral deposits, artificial fertilizers, animal waste, or organic decay (as the product of bacteria, but not bacteria specifically associated with the plant). Nitrates absorbed in this fashion are converted to nitrites by the enzyme nitrate reductase, and then converted to ammonia by another enzyme called nitrite reductase. Nitrogen compounds are basic building blocks in animal biology as well. Animals use nitrogen-containing amino acids from plant sources, as starting materials for all nitrogen-compound animal biochemistry, including the manufacture of proteins and nucleic acids. Plant-feeding insects are dependent on nitrogen in their diet, such that varying the amount of nitrogen fertilizer applied to a plant can affect the reproduction rate of insects feeding on fertilized plants.[18] Soluble nitrate is an important limiting factor in the growth of certain bacteria in ocean waters. In many places in the world, artificial fertilizers applied to crop-lands to increase yields result in run-off delivery of soluble nitrogen to oceans at river mouths. This process can result in eutrophication of the water, as nitrogen-driven bacterial growth depletes water oxygen to the point that all higher organisms die. Well-known "dead zone" areas in the U.S. Gulf Coast and the Black Sea are due to this important polluting process.
8
Nitrogen Many saltwater fish manufacture large amounts of trimethylamine oxide to protect them from the high osmotic effects of their environment (conversion of this compound to dimethylamine is responsible for the early odor in not fresh saltwater fish [19] . In animals, free radical nitric oxide (NO) (derived from an amino acid), serves as an important regulatory molecule for circulation. Animal metabolism of NO results in production of nitrite. Animal metabolism of nitrogen in proteins generally results in excretion of urea, while animal metabolism of nucleic acids results in excretion of urea and uric acid. The characteristic odor of animal flesh decay is caused by the creation of long-chain, nitrogen-containing amines, such as putrescine and cadaverine. Decay of organisms and their waste products may produce small amounts of nitrate, but most decay eventually returns nitrogen content to the atmosphere, as molecular nitrogen . The circulation of nitrogen from atmosphere to organic compounds and back is referred to as the nitrogen cycle.
Safety Rapid release of nitrogen gas into an enclosed space can displace oxygen, and therefore represents an asphyxiation hazard. This may happen with few warning symptoms, since the human carotid body is a relatively slow and a poor low-oxygen (hypoxia) sensing system.[20] An example occurred shortly before the launch of the first Space Shuttle mission in 1981, when two technicians lost consciousness and died after they walked into a space located in the Shuttle's Mobile Launcher Platform that was pressurized with pure nitrogen as a precaution against fire. The technicians would have been able to exit the room if they had experienced early symptoms from nitrogen-breathing. When inhaled at high partial pressures (more than about 4 bar, encountered at depths below about 30 m in scuba diving) nitrogen begins to act as an anesthetic agent. It can cause nitrogen narcosis, a temporary semi-anesthetized state of mental impairment similar to that caused by nitrous oxide.[21] [22] Nitrogen also dissolves in the bloodstream and body fats. Rapid decompression (particularly in the case of divers ascending too quickly, or astronauts decompressing too quickly from cabin pressure to spacesuit pressure) can lead to a potentially fatal condition called decompression sickness (formerly known as caisson sickness or more commonly, the "bends"), when nitrogen bubbles form in the bloodstream, nerves, joints, and other sensitive or vital areas.[23] [24] Other "inert" gases (those gases other than carbon dioxide and oxygen) cause the same effects from bubbles composed of them, so replacement of nitrogen in breathing gases may prevent nitrogen narcosis, but does not prevent decompression sickness.[25] Direct skin contact with liquid nitrogen will eventually cause severe frostbite (cryogenic burns). This may happen almost instantly on contact, depending on the form of liquid nitrogen. Bulk liquid nitrogen causes less rapid freezing than a spray of nitrogen mist (such as is used to freeze certain skin growths in the practice of dermatology). The extra surface area provided by nitrogen-soaked materials is also important, with soaked clothing or cotton causing far more rapid damage than a spill of direct liquid to skin. Full "contact" between naked skin and large droplets or pools of undisturbed liquid nitrogen may be prevented for a few seconds by a layer of insulating gas from the Leidenfrost effect. However, liquid nitrogen applied to skin in mists, and on fabrics, bypasses this effect.
9
Nitrogen
10
See also • • • • • • • •
Industrial gas Liquid nitrogen Nitrogen asphyxiation Nitrogenomics Nutrient Reactive nitrogen species Tetranitrogen TKN
Further reading • Garrett, Reginald H.; Grisham, Charles M. (1999). Biochemistry (2nd ed.). Fort Worth: Saunders College Publ.. ISBN 0030223180. • Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. ISBN 0080220576. • "Nitrogen [26]". Los Alamos National Laboratory. 2003-10-20. http:/ / periodic. lanl. gov/ elements/ 7. html.
External links • • • • • • • •
Etymology of Nitrogen [27] Why high nitrogen density in explosives? [28] WebElements.com – Nitrogen [29] It's Elemental – Nitrogen [30] Schenectady County Community College – Nitrogen Nitrogen N2 Properties, Uses, Applications [32] Handling procedures for liquid nitrogen [33] Material Safety Data Sheet [34]
[31]
References [1] " A new molecule and a new signature - Chemistry - tetranitrogen (http:/ / www. findarticles. com/ p/ articles/ mi_m1200/ is_7_161/ ai_83477565)". Science News. February 16, 2002. . Retrieved 2007-08-18. [2] " Polymeric nitrogen synthesized (http:/ / www. physorg. com/ news693. html)". physorg.com. 2004-08-05. . Retrieved 2009-06-22. [3] Karl Heinz Neeb (1997). The Radiochemistry of Nuclear Power Plants with Light Water Reactors. Berlin-New York: Walter de Gruyter. [4] Richard R. Schrock (2005). "Catalytic Reduction of Dinitrogen to Ammonia at a Single Molybdenum Center". Acc. Chem. Res. 38: 955–962. doi: 10.1021/ar0501121 (http:/ / dx. doi. org/ 10. 1021/ ar0501121). [5] Fryzuk, M. D. and Johnson, S. A. (2000). "The continuing story of dinitrogen activation". Coordination Chemistry Reviews 200–202: 379. doi: 10.1016/S0010-8545(00)00264-2 (http:/ / dx. doi. org/ 10. 1016/ S0010-8545(00)00264-2). [6] Croswell, Ken (February 1996). Alchemy of the Heavens (http:/ / kencroswell. com/ alchemy. html). Anchor. ISBN 0-385-47214-5. . [7] Daved M. Meyer, Jason A. Cardelli, and Ulysses J. Sofia (1997). " Abundance of Interstellar Nitrogen (http:/ / arxiv. org/ abs/ astro-ph/ 9710162v1)". arXiv. . Retrieved 2007-12-24. [8] Calvin J. Hamilton. " Titan (Saturn VI) (http:/ / www. solarviews. com/ eng/ titan. htm)". Solarviews.com. . Retrieved 2007-12-24. [9] Reich, Murray; Kapenekas, Harry (1957). "Nitrogen Purfication. Pilot Plant Removal of Oxygen". Industrial & Engineering Chemistry 49: 869. doi: 10.1021/ie50569a032 (http:/ / dx. doi. org/ 10. 1021/ ie50569a032).
Nitrogen [10] ed. by Charlie Harding ... Royal Society Chemistry; Open University. (2002). Elements of the p Block (http:/ / books. google. de/ books?id=W0HW8wgmQQsC& pg=PA90). Cambridge: Royal Society of Chemistry. ISBN 9780854046904. . [11] " Why don't they use normal air in race car tires? (http:/ / auto. howstuffworks. com/ question594. htm)". Howstuffworks. . Retrieved 2006-07-22. [12] " Diffusion, moisture and tyre expansion (http:/ / www. cartalk. com/ content/ columns/ Archive/ 1997/ September/ 05. html)". Car Talk. . Retrieved 2006-07-22. [13] " Is it better to fill your tires with nitrogen instead of air? (http:/ / www. straightdope. com/ columns/ 070216. html)". The Straight Dope. . Retrieved 2007-02-16. [14] G. J. Van Amerongen (1946). "The Permeability of Different Rubbers to Gases and Its Relation to Diffusivity and Solubility". Journal of Applied Physics 17 (11): 972–985. doi: 10.1063/1.1707667 (http:/ / dx. doi. org/ 10. 1063/ 1. 1707667). [15] Kemmochi, Y (2002). "Centrifugal concentrator for the substitution of nitrogen blow-down micro-concentration in dioxin/polychlorinated biphenyl sample preparation". Journal of Chromatography A 943: 295. doi: 10.1016/S0021-9673(01)01466-2 (http:/ / dx. doi. org/ 10. 1016/ S0021-9673(01)01466-2). [16] " How does the widget in a beer can work? (http:/ / recipes. howstuffworks. com/ question446. htm)". Howstuffworks. . [17] Rakov, Vladimir A.; Uman, Martin A. (2007). Lightning: Physics and Effects. Cambridge University Press. p. 508. ISBN 9780521035415. [18] Jahn, GC, LP Almazan, and J Pacia (2005). " Effect of nitrogen fertilizer on the intrinsic rate of increase of the rusty plum aphid, Hysteroneura setariae (Thomas) (Homoptera: Aphididae) on rice (Oryza sativa L.) (http:/ / puck. esa. catchword. org/ vl=33435372/ cl=21/ nw=1/ rpsv/ cw/ esa/ 0046225x/ v34n4/ s26/ p938)". Environmental Entomology 34 (4): 938–943. . [19] Nielsen, Mk; Jørgensen, Bm (Jun 2004). "Quantitative relationship between trimethylamine oxide aldolase activity and formaldehyde accumulation in white muscle from gadiform fish during frozen storage.". Journal of agricultural and food chemistry 52 (12): 3814–22. doi: 10.1021/jf035169l (http:/ / dx. doi. org/ 10. 1021/ jf035169l). ISSN 0021-8561 (http:/ / worldcat. org/ issn/ 0021-8561). PMID 15186102. [20] " Biology Safety - Cryogenic materials. The risks posed by them (http:/ / www. bath. ac. uk/ internal/ bio-sci/ bbsafe/ asphyx. htm)". University of Bath. . Retrieved 2007-01-03. [21] Fowler, B; Ackles, KN; Porlier, G (1985). " Effects of inert gas narcosis on behavior--a critical review. (http:/ / archive. rubicon-foundation. org/ 3019)". Undersea Biomed. Res. 12 (4): 369–402. ISSN 0093-5387 (http:/ / worldcat. org/ issn/ 0093-5387). OCLC 2068005 (http:/ / worldcat. org/ oclc/ 2068005). PMID 4082343. . Retrieved 2008-09-21. [22] W. H. Rogers; G. Moeller (1989). " Effect of brief, repeated hyperbaric exposures on susceptibility to nitrogen narcosis (http:/ / archive. rubicon-foundation. org/ 2522)". Undersea Biomed. Res. 16 (3): 227–32. ISSN 0093-5387 (http:/ / worldcat. org/ issn/ 0093-5387). OCLC 2068005 (http:/ / worldcat. org/ oclc/ 2068005). PMID 2741255. . Retrieved 2008-09-21. [23] Acott, C. (1999). " A brief history of diving and decompression illness. (http:/ / archive. rubicon-foundation. org/ 6004)". South Pacific Underwater Medicine Society journal 29 (2). ISSN 0813-1988 (http:/ / worldcat. org/ issn/ 0813-1988). OCLC 16986801 (http:/ / worldcat. org/ oclc/ 16986801). . Retrieved 2008-09-21. [24] Kindwall, E. P.; A. Baz; E. N. Lightfoot; E. H. Lanphier; A. Seireg. (1975). " Nitrogen elimination in man during decompression. (http:/ / archive. rubicon-foundation. org/ 2741)". Undersea Biomed. Res. 2 (4): 285–97. ISSN 0093-5387 (http:/ / worldcat. org/ issn/ 0093-5387). OCLC 2068005 (http:/ / worldcat. org/ oclc/ 2068005). PMID 1226586. . Retrieved 2008-09-21. [25] US Navy Diving Manual, 6th revision (http:/ / www. supsalv. org/ 00c3_publications. asp?destPage=00c3& pageID=3. 9). United States: US Naval Sea Systems Command. 2006. . Retrieved 2008-04-24. [26] [27] [28] [29] [30] [31] [32] [33] [34]
http:/ / periodic. lanl. gov/ elements/ 7. html http:/ / www. balashon. com/ 2008/ 07/ neter-and-nitrogen. html http:/ / www. newton. dep. anl. gov/ askasci/ chem99/ chem99306. htm http:/ / www. webelements. com/ nitrogen/ http:/ / education. jlab. org/ itselemental/ ele007. html http:/ / www. sunysccc. edu/ academic/ mst/ ptable/ n. html http:/ / www. uigi. com/ nitrogen. html http:/ / www. 2spi. com/ catalog/ instruments/ nitrodew-supp. html http:/ / www. safety. vanderbilt. edu/ pdf/ hcs_msds/ NitrogenCryo_G103_06_04. pdf
11
Article Sources and Contributors
Article Sources and Contributors Nitrogen Source: http://en.wikipedia.org/w/index.php?oldid=305282750 Contributors: 1243trel;ghsfdk;lgsfd, 1297, 130.225.29.xxx, 8472, A3r0, AAA!, Adashiel, Addshore, Adrian.benko, Agamemnon2, Ahoerstemeier, Ajaxkroon, AlainV, Alan012, Alansohn, Alexfusco5, AlexiusHoratius, Alexwcovington, All Is One, AlphaEta, Alphachimp, Andre Engels, AndreasJS, AndreniW, Andres, AnonGuy, Anonymous Dissident, Antandrus, Anthere, Anthony Appleyard, Apothecia, Arakunem, Archer3, Archie Paulson, Archimerged, ArielGold, Arjun01, Asldkjaskdjsalkjdk, Atlant, Atulsnischal, Avant Guard, Awsomesoccerman, Axemanstan, Axlenz, BRG, Badocter, Bayerischermann, Bayou Banjo, Beetstra, Beland, Benbest, Bender235, Benjah-bmm27, Bert Hickman, Bfesser, Bgdsnezana, Biochemnick, Bkell, BlueEarth, Bobianite, Bobo The Ninja, Bobo192, Bogey97, Bonadea, BrianGV, Brianga, BrockF5, BrokenSegue, BrokenSphere, Bsgreader, Bucketsofg, Bushcarrot, C0nanPayne, CYD, Cabazon, Caesura, Calabraxthis, Calliopejen1, Camw, Can't sleep, clown will eat me, Canada rocks06, CanadianLinuxUser, Cando12, CanisRufus, CapitalR, Capricorn42, Captain panda, Carnildo, Castaneda, CatherineMunro, Causesobad, Cdc, Ceranthor, Cflm001, CharlotteWebb, ChemNerd, ChicXulub, Chris Dybala, Chris the speller, ChrisSteinbach, Chrisjj, Chrislk02, Chriswiki, Chun-hian, CiTrusD, Cimon Avaro, ClockworkSoul, Closedmouth, Cobaltbluetony, Colbuckshot, Cometstyles, Controlfreeke, Conversion script, Coppertwig, Correogsk, Cosmium, Crazynas, Crazysane, Cremepuff222, Ctjf83, D, DMacks, DVD R W, Da monster under your bed, DaDrought3, Dale Arnett, DancingPenguin, Dar-Ape, DarkAudit, Darkspots, Darrien, Darth Panda, David D., David Gerard, David Latapie, David Schaich, David.Monniaux, DavidLevinson, Dbtfz, Dcandeto, Deglr6328, Delldot, Delta G, Dendodge, Denelson83, Deor, DerHexer, Dethme0w, Deville, Devonmcdonough, Dgrant, Discospinster, DoktorDec, Donald Albury, DoubleBlue, Doulos Christos, Drini, Drmaik, Dryguy, Dwmyers, Dycedarg, Dysepsion, EL Willy, ESkog, EarthPerson, Eddpayne, Edgar181, Edgjerp, Edsanville, Eeekster, Effendi13, Egil, Ekton, El C, ElAmericano, Eldin raigmore, Eleassar, Elkman, Em Mitchell, Emmelie, Emperorbma, Eng02019, Enok Walker, Epbr123, Eric119, Erik Zachte, EstebanF, Esteffect, Evercat, Everyking, Ewlyahoocom, Excirial, Fatass1220, Femto, Fermion, Finn-Zoltan, Firsfron, FiveTurkey, Flewis, Flowerpotman, Foobar, FrancoGG, Frank Lofaro Jr., FrankTobia, Freakmighty, Free to rhyme, FreplySpang, Fresheneesz, Furrykef, GT, Gail, Gaius Cornelius, Gamargamar, Gamedevil, Gen.Sec., Gene Hobbs, Gene Nygaard, Gibo333, Giftlite, Gittinsj, Gjahn, Gkenda, Glenn, Glossyn, Gman124, Goalie94, Gogo Dodo, GraemeL, Granf, Grasys, Grendelkhan, Griffinofwales, Grunt, Gtg204y, Gurch, Gurchzilla, Guyinypsi, Gwernol, Habsfan12337, Hadal, Hak-kâ-ngìn, HarryAlffa, Hdt83, Heightwatcher, Heimstern, Hellbus, Henry W. 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12
Oxygen
1
Oxygen nitrogen ← oxygen → fluorine ↑ O ↓ S
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
8O Periodic table
Appearance liquid oxygen with bubbles of oxygen gas
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensityMelting pointBoiling pointCritical pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa at T/K
1
10
100
1k
10 k
100 k
61
73
90
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 3388.3 kJ·mol−1 3rd: 5300.5 kJ·mol−1Covalent radiusVan der Waals radius Miscellaneous Crystal structureMagnetic orderingThermal conductivitySpeed of soundCAS registry number Most stable isotopes Main article: Isotopes of oxygen
Oxygen
iso
2
N.A.
half-life
16
99.76%
16
17
0.039%
17
18
0.201%
18
O O O
DM
DE (MeV)
DP
O is stable with 8 neutron O is stable with 9 neutron O is stable with 10 neutron
oxygen, O, 8 nonmetal, chalcogens 16, 2, p15.9994(3) g·mol−1 1s2 2s2 2p4 2, 6 (Image) gas (0 °C, 101.325 kPa) 1.429 g/L 54.36 K,-218.79 °C,-361.82 °F 90.20 K,-182.95 °C,-297.31 °F 154.59 K, 5.043 MPa (O2) 0.444 kJ·mol−1 (O2) 6.82 kJ·mol−1 (25 °C) (O2) 29.378 J·mol−1·K−12, 1, −1, −2 (neutral oxide) 3.44 (Pauling scale) 1st: 1313.9 kJ·mol−166±2 pm 152 pm cubic paramagnetic (300 K) 26.58x10-3 W·m−1·K−1 (gas, 27 °C) 330 m/s 7782-44-7 Oxygen (pronounced /ˈɒksɨdʒɨn/, from the Greek roots ὀξύς (oxys) (acid, literally "sharp", from the taste of acids) and -γενής (-genēs) (producer, literally begetter) is the element with atomic number 8 and represented by the symbol O. It is a member of the chalcogen group on the periodic table, and is a highly reactive nonmetallic period 2 element that readily forms compounds (notably oxides) with almost all other elements. At standard temperature and pressure two atoms of the element bind to form dioxygen, a colorless, odorless, tasteless diatomic gas with the formula O2. Oxygen is the third most abundant element in the universe by mass after hydrogen and helium[1] and the most abundant element by mass in the Earth's crust.[2] Diatomic oxygen gas constitutes 20.9% of the volume of air.[3] All major classes of structural molecules in living organisms, such as proteins, carbohydrates, and fats, contain oxygen, as do the major inorganic compounds that comprise animal shells, teeth, and bone. Oxygen in the form of O2 is produced from water by cyanobacteria, algae and plants during photosynthesis and is used in cellular respiration for all complex life. Oxygen is toxic to obligately anaerobic organisms, which were the dominant form of early life on Earth until O2 began to accumulate in the atmosphere 2.5 billion years ago.[4] Another form (allotrope) of oxygen, ozone (O3), helps protect the biosphere from ultraviolet radiation with the high-altitude ozone layer, but is a pollutant near the surface where it is a by-product of smog. At even higher low earth orbit altitudes atomic oxygen is a significant presence and a cause of erosion for spacecraft.[5] Oxygen was independently discovered by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, and Joseph Priestley in Wiltshire, in 1774, but Priestley is often given priority because his publication came out in print first. The name oxygen was coined in 1777 by Antoine Lavoisier,[6] whose experiments with oxygen helped to discredit the then-popular phlogiston theory of combustion and corrosion. Oxygen is produced industrially by fractional distillation of liquefied air, use of zeolites to remove carbon dioxide and nitrogen from air, electrolysis of water and other means. Uses of oxygen include the production of steel, plastics and textiles; rocket propellant; oxygen therapy; and life support in aircraft, submarines, spaceflight and diving.
Oxygen
Characteristics Structure At standard temperature and pressure, oxygen is a colorless, odorless gas with the molecular formula O2, in which the two oxygen atoms are chemically bonded to each other with a spin triplet electron configuration. This bond has a bond order of two, and is often simplified in description as a double bond[7] or as a combination of one two-electron bond and two three-electron bonds.[8] Triplet oxygen is the ground state of the O2 molecule.[9] The electron configuration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals.[10] These orbitals are classified as antibonding (weakening the bond order from three to two), so the diatomic oxygen bond is weaker than the diatomic nitrogen triple bond in which all bonding molecular orbitals are filled, but some antibonding orbitals are not.[9] In normal triplet form, O2 molecules are paramagnetic—they form a magnet in the presence of a magnetic field—because of the spin magnetic moments of the unpaired electrons in the molecule, and the negative exchange energy between neighboring O2 molecules.[11] Liquid oxygen is attracted to a magnet to a sufficient extent that, in laboratory demonstrations, a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet.[12] [13] Singlet oxygen, a name given to several higher-energy species of molecular O2 in which all the electron spins are paired, is much more reactive towards common organic molecules. In nature, singlet oxygen is commonly formed from water during photosynthesis, using the energy of sunlight.[14] It is also produced in the troposphere by the photolysis of ozone by light of short wavelength,[15] and by the immune system as a source of active oxygen.[16] Carotenoids in photosynthetic organisms (and possibly also in animals) play a major role in absorbing energy from singlet oxygen and converting it to the unexcited ground state before it can cause harm to tissues.[17]
3
Oxygen
4
Allotropes The common allotrope of elemental oxygen on Earth is called dioxygen, O2. It has a bond length of 121 pm and a bond energy of 498 kJ·mol-1.[18] This is the form that is used by complex forms of life, such as animals, in cellular respiration (see Biological role) and is the form that is a major part of the Earth's atmosphere (see Occurrence). Other aspects of O2 are covered in the remainder of this article. Trioxygen (O3) is usually known as ozone and is a very reactive allotrope of oxygen that is damaging to lung tissue.[19] Ozone is produced in the upper atmosphere when O2 combines with atomic oxygen made by the splitting of O2 by ultraviolet (UV) radiation.[6] Since ozone absorbs strongly in the UV region of the spectrum, the ozone layer of the upper atmosphere functions as a protective radiation shield for the planet.[6] Near the Earth's surface, however, it is a pollutant formed as a by-product of automobile exhaust.[20]
Ozone is a rare gas on Earth found mostly in the stratosphere
The metastable molecule tetraoxygen (O4) was discovered in 2001,[21] [22] and was assumed to exist in one of the six phases of solid oxygen. It was proven in 2006 that this phase, created by pressurizing O2 to 20 GPa, is in fact a rhombohedral O8 cluster.[23] This cluster has the potential to be a much more powerful oxidizer than either O2 or O3 and may therefore be used in rocket fuel.[21] [22] A metallic phase was discovered in 1990 when solid oxygen is subjected to a pressure of above 96 GPa[24] and it was shown in 1998 that at very low temperatures, this phase becomes superconducting.[25]
Physical properties Oxygen is more soluble in water than nitrogen; water contains approximately 1 molecule of O2 for every 2 molecules of N2, compared to an atmospheric ratio of approximately 1:4. The solubility of oxygen in water is temperature-dependent, and about twice as much (14.6 mg·L−1) dissolves at 0 °C than at 20 °C (7.6 mg·L−1).[26] [27] At 25 °C and 1 standard atmosphere (101.3 kPa) of air, freshwater contains about 6.04 milliliters (mL) of oxygen per liter, whereas seawater contains about 4.95 mL per liter.[28] At 5 °C the solubility increases to 9.0 mL (50% more than at 25 °C) per liter for water and 7.2 mL (45% more) per liter for sea water. Oxygen condenses at 90.20 K (−182.95 °C, −297.31 °F), and freezes at 54.36 K (−218.79 °C, −361.82 °F).[29] Both liquid and solid O2 are clear substances with a light sky-blue color caused by absorption in the red (in contrast with the blue color of the sky, which is due to Rayleigh scattering of blue light). High-purity liquid O2 is usually obtained by the fractional distillation of liquefied air;[30] Liquid oxygen may also be produced by condensation out of air, using liquid nitrogen as a coolant. It is a highly reactive substance and must be segregated from combustible materials.[31]
Oxygen
5
Isotopes and stellar origin Naturally occurring oxygen is composed of three stable isotopes, 16O, 17O, and 18O, with 16O being the most abundant (99.762% natural abundance).[32] Most 16O is synthesized at the end of the helium fusion process in stars but some is made in the neon burning process.[33] 17O is primarily made by the burning of hydrogen into helium during the CNO cycle, making it a common isotope in the hydrogen burning zones of stars.[33] Most 18O is produced when 14N (made abundant from CNO burning) captures a 4He nucleus, making 18O common in the helium-rich zones of stars.[33]
Late in a massive star's life, 16O concentrates in the O-shell, 17O in the H-shell and 18O in the He-shell
Fourteen radioisotopes have been characterized, the 15 most stable being O with a half-life of 14 122.24 seconds (s) and O with a half-life of 70.606 s.[32] All of the remaining radioactive isotopes have half-lives that are less than 27 s and the majority of these have half-lives that are less than 83 milliseconds.[32] The most common decay mode of the isotopes lighter than 16 O is β+ decay[34] [35] [36] to yield nitrogen, and the most common mode for the isotopes heavier than 18O is beta decay to yield fluorine.[32]
Occurrence Oxygen is the most abundant chemical element, by mass, in our biosphere, air, sea and land. Oxygen is the third most abundant chemical element in the universe, after hydrogen and helium.[1] About 0.9% of the Sun's mass is oxygen.[3] Oxygen constitutes 49.2% of the Earth's crust by mass[2] and is the major component of the world's oceans (88.8% by mass).[3] Oxygen gas is the second most common component of the Earth's atmosphere, taking up 21.0% of its volume and 23.1% of its mass (some 1015 tonnes).[3] [37] [38] Earth is unusual among the planets of the Solar System in having such a high concentration of oxygen gas in its atmosphere: Mars (with 0.1% O2 by volume) and Venus have far lower concentrations. However, the O2 surrounding these other planets is produced solely by ultraviolet radiation impacting oxygen-containing molecules such as carbon dioxide. The unusually high concentration of oxygen gas on Earth is the result of the oxygen cycle. This biogeochemical cycle describes the movement of oxygen within and between its three main reservoirs on Earth: the atmosphere, the biosphere, and the lithosphere. The main driving factor of the oxygen cycle is photosynthesis, which is responsible for modern Earth's atmosphere. Photosynthesis releases oxygen into the atmosphere, while respiration and decay Cold water holds more dissolved O2. remove it from the atmosphere. In the present equilibrium, production and consumption occur at the same rate of roughly 1/2000th of the entire atmospheric oxygen per year.
Oxygen
6
Free oxygen also occurs in solution in the world's water bodies. The increased solubility of O2 at lower temperatures (see Physical properties) has important implications for ocean life, as polar oceans support a much higher density of life due to their higher oxygen content.[39] Polluted water may have reduced amounts of O2 in it, depleted by decaying algae and other biomaterials (see eutrophication). Scientists assess this aspect of water quality by measuring the water's biochemical oxygen demand, or the amount of O2 needed to restore it to a normal concentration.[40]
Biological role Photosynthesis and respiration In nature, free oxygen is produced by the light-driven splitting of water during oxygenic photosynthesis. Green algae and cyanobacteria in marine environments provide about 70% of the free oxygen produced on earth and the rest is produced by terrestrial plants.[41] A simplified overall formula for photosynthesis is:[42] 6 CO2 + 6 H2O + photons → C6H12O6 + 6 O2 (or simply carbon dioxide + water + sunlight → glucose + dioxygen)
Photolytic oxygen evolution occurs in the thylakoid membranes of photosynthetic organisms and requires the energy of four photons.[43] Many steps are involved, but the result is the formation of a proton gradient across the thylakoid membrane, which is used to synthesize ATP via photophosphorylation.[44] The O2 remaining after oxidation of the water molecule is released into the atmosphere.[45]
Photosynthesis splits water to liberate O2 and fixes CO2 into sugar
Oxygen
7 Molecular dioxygen, O2, is essential for cellular respiration in all aerobic organisms. Oxygen is used in mitochondria to help generate adenosine triphosphate (ATP) during oxidative phosphorylation. The reaction for aerobic respiration is essentially the reverse of photosynthesis and is simplified as: C6H12O6 + 6 O2 → 6 CO2 + 6 H2O + 2880 kJ·mol-1
Relation between photosynthesis and respiration. Oxygen (at left) is consumed in respiration of organic compounds to form carbon dioxide and water. These can again produce oxygen and organic compounds in photosynthesis.
In vertebrates, O2 is diffused through membranes in the lungs and into red blood cells. Hemoglobin binds O2, changing its color from bluish red to bright red.[19] [46] Other animals use hemocyanin (molluscs and some arthropods) or hemerythrin (spiders and lobsters).[37] A liter of blood can dissolve 200 cm3 of O2.[37]
Reactive oxygen species, such as superoxide ion (O−2) and hydrogen peroxide (H2O2), are dangerous by-products of oxygen use in organisms.[37] Parts of the immune system of higher organisms, however, create peroxide, superoxide, and singlet oxygen to destroy invading microbes. Reactive oxygen species also play an important role in the hypersensitive response of plants against pathogen attack.[44] An adult human in rest inhales 1.8 to 2.4 grams of oxygen per minute.[47] This amounts to more than 6 billion tonnes of oxygen inhaled by humanity per year. [48]
Build-up in the atmosphere Free oxygen gas was almost nonexistent in Earth's atmosphere before photosynthetic archaea and bacteria evolved. Free oxygen first appeared in significant quantities during the Paleoproterozoic era (between 2.5 and 1.6 billion years ago). At first, the oxygen combined with dissolved iron in the oceans to form banded iron formations. Free oxygen started to gas out of the oceans 2.7 billion years ago, reaching 10% of its present level around 1.7 billion years ago.[49]
O2 build-up in Earth's atmosphere: 1) no O2 produced; 2) O2 produced, but absorbed in oceans & seabed rock; 3) O2 starts to gas out of the oceans, but is absorbed by land surfaces and formation of ozone layer; 4-5) O2 sinks filled and the gas accumulates
The presence of large amounts of dissolved and free oxygen in the oceans and atmosphere may have driven most of the anaerobic organisms then living to extinction during the
Oxygen
8
oxygen catastrophe about 2.4 billion years ago. However, cellular respiration using O2 enables aerobic organisms to produce much more ATP than anaerobic organisms, helping the former to dominate Earth's biosphere.[50] Photosynthesis and cellular respiration of O2 allowed for the evolution of eukaryotic cells and ultimately complex multicellular organisms such as plants and animals. Since the beginning of the Cambrian era 540 million years ago, O2 levels have fluctuated between 15% and 30% by volume.[51] Towards the end of the Carboniferous era (about 300 million years ago) atmospheric O2 levels reached a maximum of 35% by volume,[51] allowing insects and amphibians to grow much larger than today's species. Human activities, including the burning of 7 billion tonnes of fossil fuels each year have had very little effect on the amount of free oxygen in the atmosphere.[11] At the current rate of photosynthesis it would take about 2,000 years to regenerate the entire O2 in the present atmosphere.[52]
History Early experiments One of the first known experiments on the relationship between combustion and air was conducted by the second century BCE Greek writer on mechanics, Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle and surrounding the vessel's neck with water resulted in some water rising into the neck.[53] Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries later Leonardo da Vinci built on Philo's work by observing that a portion of air is consumed during combustion and respiration.[54] In the late 17th century, Robert Boyle proved that air is necessary for combustion. English chemist John Mayow refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus or just nitroaereus.[55] In one experiment he found that placing either a mouse or a lit candle in a closed container over water caused the water to rise and replace one-fourteenth of the air's volume before extinguishing the subjects.[56] From this he surmised that nitroaereus is consumed in both respiration and combustion.
Philo's experiment inspired later investigators
Mayow observed that antimony increased in weight when heated, and inferred that the nitroaereus must have combined with it.[55] He also thought that the lungs separate nitroaereus from air and pass it into the blood and that animal heat and muscle movement result from the reaction of nitroaereus with certain substances in the body.[55] Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione".[56]
Oxygen
9
Phlogiston theory Robert Hooke, Ole Borch, Mikhail Lomonosov, and Pierre Bayen all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as an element.[26] This may have been in part due to the prevalence of the philosophy of combustion and corrosion called the phlogiston theory, which was then the favored explanation of those processes.
Stahl helped develop and popularize the phlogiston theory.
Established in 1667 by the German alchemist J. J. Becher, and modified by the chemist Georg Ernst Stahl by 1731,[57] phlogiston theory stated that all combustible materials were made of two parts. One part, called phlogiston, was given off when the substance containing it was burned, while the dephlogisticated part was thought to be its true form, or calx.[54]
Highly combustible materials that leave little residue, such as wood or coal, were thought to be made mostly of phlogiston; whereas non-combustible substances that corrode, such as iron, contained very little. Air did not play a role in phlogiston theory, nor were any initial quantitative experiments conducted to test the idea; instead, it was based on observations of what happens when something burns, that most common objects appear to become lighter and seem to lose something in the process.[54] The fact that a substance like wood actually gains overall weight in burning was hidden by the buoyancy of the gaseous combustion products. Indeed one of the first clues that the phlogiston theory was incorrect was that metals, too, gain weight in rusting (when they were supposedly losing phlogiston).
Discovery Oxygen was first discovered by Swedish pharmacist Carl Wilhelm Scheele. He had produced oxygen gas by heating mercuric oxide and various nitrates by about 1772.[3] [54] Scheele called the gas 'fire air' because it was the only known supporter of combustion. He wrote an account of this discovery in a manuscript he titled Treatise on Air and Fire, which he sent to his publisher in 1775. However, that document was not published until 1777.[58] Carl Wilhelm Scheele beat Priestley to the discovery but published afterwards.
Oxygen
Joseph Priestley is usually given priority in the discovery
10 In the meantime, an experiment was conducted by the British clergyman Joseph Priestley on August 1, 1774 focused sunlight on mercuric oxide (HgO) inside a glass tube, which liberated a gas he named 'dephlogisticated air'.[3] He noted that candles burned brighter in the gas and that a mouse was more active and lived longer while breathing it. After breathing the gas himself, he wrote: "The feeling of it to my lungs was not sensibly different from that of common air, but I fancied that my breast felt peculiarly light and easy for some time afterwards."[26] Priestley published his findings in 1775 in a paper titled "An Account of Further Discoveries in Air" which was included in the second volume of his book titled Experiments and Observations on Different Kinds of Air.[54] [59] Because he published his findings first, Priestley is usually given priority in the discovery.
The noted French chemist Antoine Laurent Lavoisier later claimed to have discovered the new substance independently. However, Priestley visited Lavoisier in October 1774 and told him about his experiment and how he liberated the new gas. Scheele also posted a letter to Lavoisier on September 30, 1774 that described his own discovery of the previously unknown substance, but Lavoisier never acknowledged receiving it (a copy of the letter was found in Scheele's belongings after his death).[58]
Lavoisier's contribution What Lavoisier did indisputably do (although this was disputed at the time) was to conduct the first adequate quantitative experiments on oxidation and give the first correct explanation of how combustion works.[3] He used these and similar experiments, all started in 1774, to discredit the phlogiston theory and to prove that the substance discovered by Priestley and Scheele was a chemical element. In one experiment, Lavoisier observed that there was no overall increase in weight when tin and air were heated in a closed container.[3] He noted that air rushed in when he opened the container, which indicated that part of the trapped air had been consumed. He also noted that the tin had increased in weight and that increase was the same as the weight of the air that rushed back in. This and other experiments on combustion were documented in his book Sur la combustion en général, which was published in 1777.[3] In that work, he proved that air is a mixture of two gases; 'vital air', which is essential to Antoine Lavoisier discredited the Phlogiston theory combustion and respiration, and azote (Gk. ἄζωτον "lifeless"), which did not support either. Azote later became nitrogen in English, although it has kept the name in French and several other European languages.[3] Lavoisier renamed 'vital air' to oxygène in 1777 from the Greek roots ὀξύς (oxys) (acid, literally "sharp," from the taste of acids) and -γενής (-genēs) (producer, literally begetter), because he mistakenly believed that oxygen was a constituent of all acids.[6] Chemists eventually determined that Lavoisier was wrong in this regard, but by that time it was too
Oxygen
11
late, the name had taken. Actually, the gas that could more appropriately have been given the description, "acid producer," is hydrogen. Oxygen entered the English language despite opposition by English scientists and the fact that the Englishman Priestley had first isolated the gas and written about it. This is partly due to a poem praising the gas titled "Oxygen" in the popular book The Botanic Garden (1791) by Erasmus Darwin, grandfather of Charles Darwin.[58]
Later history John Dalton's original atomic hypothesis assumed that all elements were monoatomic and that the atoms in compounds would normally have the simplest atomic ratios with respect to one another. For example, Dalton assumed that water's formula was HO, giving the atomic mass of oxygen as 8 times that of hydrogen, instead of the modern value of about 16.[60] In 1805, Joseph Louis Gay-Lussac and Alexander von Humboldt showed that water is formed of two volumes of hydrogen and one volume of oxygen; and by 1811 Amedeo Avogadro had arrived at the correct interpretation of water's composition, based on what is now called Avogadro's law and the assumption of diatomic elemental molecules.[61] [62] By the late 19th century scientists realized that air
Robert H. Goddard and a liquid oxygen-gasoline rocket
could be liquefied, and its components isolated, by compressing and cooling it. Using a cascade method, Swiss chemist and physicist Raoul Pierre Pictet evaporated liquid sulfur dioxide in order to liquefy carbon dioxide, which in turn was evaporated to cool oxygen gas enough to liquefy it. He sent a telegram on December 22, 1877 to the French Academy of Sciences in Paris announcing his discovery of liquid oxygen.[63] Just two days later, French physicist Louis Paul Cailletet announced his own method of liquefying molecular oxygen.[63] Only a few drops of the liquid were produced in either case so no meaningful analysis could be conducted. Oxygen was liquified in stable state for the first time on March 29, 1877 by Polish scientists from Jagiellonian University, Zygmunt Wróblewski and Karol Olszewski.[64] In 1891 Scottish chemist James Dewar was able to produce enough liquid oxygen to study.[11] The first commercially viable process for producing liquid oxygen was independently developed in 1895 by German engineer Carl von Linde and British engineer William Hampson. Both men lowered the temperature of air until it liquefied and then distilled the component gases by boiling them off one at a time and capturing them.[65] Later, in 1901, oxyacetylene welding was demonstrated for the first time by burning a mixture of acetylene and compressed O2. This method of welding and cutting metal later became common.[65] In 1923 the American scientist Robert H. Goddard became the first person to develop a rocket engine; the engine used gasoline for fuel and liquid oxygen as the oxidizer. Goddard successfully flew a small liquid-fueled rocket 56 m at 97 km/h on March 16, 1926 in Auburn, Massachusetts, USA.[65] [66]
Oxygen
12
Industrial production Two major methods are employed to produce 100 million tonnes of O2 extracted from air for industrial uses annually.[58] The most common method is to fractionally distill liquefied air into its various components, with nitrogen N2 distilling as a vapor while oxygen O2 is left as a liquid.[58] The other major method of producing O2 gas involves passing a stream of clean, dry air through one bed of a pair of identical zeolite molecular sieves, which absorbs the nitrogen and delivers a gas stream that is 90% to 93% O2.[58] Simultaneously, nitrogen gas is released from the other nitrogen-saturated zeolite bed, by reducing the chamber operating pressure and diverting part of the oxygen gas from the producer bed through it, in the reverse direction of flow. After a set cycle time the operation of the two beds is interchanged, thereby allowing for a continuous supply of gaseous oxygen to be pumped through a pipeline. This is known as pressure swing adsorption. Oxygen gas is increasingly obtained by these non-cryogenic technologies (see also the related vacuum swing adsorption).[67] Oxygen gas can also be produced through electrolysis of water into molecular oxygen and hydrogen. DC electricity must be used: if AC is used, the gases in each limb consist of hydrogen and oxygen in the explosive ratio 2:1. Contrary to popular belief, the 2:1 ratio observed in the DC electrolysis of acidified water does not prove that the empirical formula of water is H2O unless certain assumptions are made about the molecular formulae of hydrogen and oxygen themselves. Hoffman electrolysis apparatus used in electrolysis of water.
A similar method is the electrocatalytic O2 evolution from oxides and oxoacids. Chemical catalysts can be used as well, such as in chemical oxygen generators or oxygen candles that are used as part of the life-support equipment on submarines, and are still part of standard equipment on commercial airliners in case of depressurization emergencies. Another air separation technology involves forcing air to dissolve through ceramic membranes based on zirconium dioxide by either high pressure or an electric current, to produce nearly pure O2 gas.[40] In large quantities, the price of liquid oxygen in 2001 was approximately $0.21/kg.[68] Since the primary cost of production is the energy cost of liquefying the air, the production cost will change as energy cost varies. For reasons of economy, oxygen is often transported in bulk as a liquid in specially insulated tankers, since one litre of liquefied oxygen is equivalent to 840 liters of gaseous oxygen at atmospheric pressure and 20 °C.[58] Such tankers are used to refill bulk liquid oxygen storage containers, which stand outside hospitals and other institutions with a need for large volumes of pure oxygen gas. Liquid oxygen is passed through heat exchangers, which convert the cryogenic liquid into gas before it enters the building. Oxygen is also stored and shipped in smaller cylinders containing the compressed gas; a form that is
Oxygen
13
useful in certain portable medical applications and oxy-fuel welding and cutting.[58]
Applications Medical Uptake of O2 from the air is the essential purpose of respiration, so oxygen supplementation is used in medicine. Oxygen therapy is used to treat emphysema, pneumonia, some heart disorders, and any disease that impairs the body's ability to take up and use gaseous oxygen.[69] Treatments are flexible enough to be used in hospitals, the patient's home, or increasingly by portable devices. Oxygen tents were once commonly used in oxygen supplementation, but have since been replaced mostly by the use of oxygen masks or nasal cannulas.[70] Hyperbaric (high-pressure) medicine uses special oxygen An oxygen concentrator in an chambers to increase the partial pressure of O2 around the [71] emphysema patient's house patient and, when needed, the medical staff. Carbon monoxide poisoning, gas gangrene, and decompression sickness (the 'bends') are sometimes treated using these devices.[72] Increased O2 concentration in the lungs helps to displace carbon monoxide from the heme group of hemoglobin.[73] [74] Oxygen gas is poisonous to the anaerobic bacteria that cause gas gangrene, so increasing its partial pressure helps kill them.[75] [76] Decompression sickness occurs in divers who decompress too quickly after a dive, resulting in bubbles of inert gas, mostly nitrogen and helium, forming in their blood. Increasing the pressure of O2 as soon as possible is part of the treatment.[69] [77] [78] Oxygen is also used medically for patients who require mechanical ventilation, often at concentrations above 21% found in ambient air.
Life support and recreational use A notable application of O2 as a low-pressure breathing gas is in modern space suits, which surround their occupant's body with pressurized air. These devices use nearly pure oxygen at about one third normal pressure, resulting in a normal blood partial pressure of O2.[79] [80] This trade-off of higher oxygen concentration for lower pressure is needed to maintain flexible spacesuits. Scuba divers and submariners also rely on artificially delivered O2, but most often use normal pressure, and/or mixtures of oxygen and air. Pure or nearly pure O2 use in diving at higher-than-sea-level pressures is usually limited to rebreather, decompression, or
Low pressure pure O2 is used in space suits
Oxygen emergency treatment use at relatively shallow depths (~6 meters depth, or less).[81] [82] Deeper diving requires significant dilution of O2 with other gases, such as nitrogen or helium, to help prevent oxygen toxicity.[81] People who climb mountains or fly in non-pressurized fixed-wing aircraft sometimes have supplemental O2 supplies.[83] Passengers traveling in (pressurized) commercial airplanes have an emergency supply of O2 automatically supplied to them in case of cabin depressurization. Sudden cabin pressure loss activates chemical oxygen generators above each seat, causing oxygen masks to drop and forcing iron filings into the sodium chlorate inside the canister.[40] A steady stream of oxygen gas is produced by the exothermic reaction. However, even this may pose a danger if inappropriately triggered: a ValuJet airplane crashed after use-date-expired O2 canisters, which were being shipped in the cargo hold, activated and caused fire. The canisters were mis-labeled as empty, and carried against dangerous goods regulations.[84] Oxygen, as a supposed mild euphoric, has a history of recreational use in oxygen bars and in sports. Oxygen bars are establishments, found in Japan, California, and Las Vegas, Nevada since the late 1990s that offer higher than normal O2 exposure for a fee.[85] Professional athletes, especially in American football, also sometimes go off field between plays to wear oxygen masks in order to get a supposed "boost" in performance. However, the reality of a pharmacological effect is doubtful; a placebo or psychological boost being the most plausible explanation.[85] Available studies support a performance boost from enriched O2 mixtures only if they are breathed during actual aerobic exercise.[86] Other recreational uses include pyrotechnic applications, such as George Goble's five-second ignition of barbecue grills.[87]
Industrial Smelting of iron ore into steel consumes 55% of commercially produced oxygen.[40] In this process, O2 is injected through a high-pressure lance into molten iron, which removes sulfur impurities and excess carbon as the respective oxides, SO2 and CO2. The reactions are exothermic, so the temperature increases to [40] 1,700 °C. Another 25% of commercially produced oxygen is used by the chemical industry.[40] Ethylene is reacted with Most commercially produced O2 is O2 to create ethylene oxide, which, in turn, is converted used to smelt iron into steel. into ethylene glycol; the primary feeder material used to manufacture a host of products, including antifreeze and polyester polymers (the precursors of many plastics and fabrics).[40] Most of the remaining 20% of commercially produced oxygen is used in medical applications, metal cutting and welding, as an oxidizer in rocket fuel, and in water treatment.[40] Oxygen is used in oxyacetylene welding burning acetylene with O2 to produce a very hot flame. In this process, metal up to 60 cm thick is first heated with a small oxy-acetylene flame and then quickly cut by a large stream of O2.[88] Rocket propulsion requires a fuel and an oxidizer. Larger rockets use liquid oxygen as their oxidizer, which is mixed and ignited with the fuel for propulsion.
14
Oxygen
Scientific Paleoclimatologists measure the ratio of oxygen-18 and oxygen-16 in the shells and skeletons of marine organisms to determine what the climate was like millions of years ago (see oxygen isotope ratio cycle). Seawater molecules that contain the lighter isotope, oxygen-16, evaporate at a slightly faster rate than water molecules containing the 12% heavier oxygen-18; this disparity increases at lower temperatures.[89] During periods of lower global 500 million years of climate change vs 18O temperatures, snow and rain from that evaporated water tends to be higher in oxygen-16, and the seawater left behind tends to be higher in oxygen-18. Marine organisms then incorporate more oxygen-18 into their skeletons and shells than they would in a warmer climate.[89] Paleoclimatologists also directly measure this ratio in the water molecules of ice core samples that are up to several hundreds of thousands of years old. Planetary geologists have measured different abundances of oxygen isotopes in samples from the Earth, the Moon, Mars, and meteorites, but were long unable to obtain reference values for the isotope ratios in the Sun, believed to be the same as those of the primordial solar nebula. However, analysis of a silicon wafer exposed to the solar wind in space and returned by the crashed Genesis spacecraft has shown that the Sun has a higher proportion of oxygen-16 than does the Earth. The measurement implies that an unknown process depleted oxygen-16 from the Sun's disk of protoplanetary material prior to the coalescence of dust grains that formed the Earth.[90] Oxygen presents two spectrophotometric absorption bands peaking at the wavelengths 687 and 760 nm. Some remote sensing scientists have proposed using the measurement of the radiance coming from vegetation canopies in those bands to characterize plant health status from a satellite platform.[91] This approach exploits the fact that in those bands it is possible to discriminate the vegetation's reflectance from its fluorescence, which is much weaker. The measurement is technically difficult owing to the low signal-to-noise ratio and the physical structure of vegetation; but it has been proposed as a possible method of monitoring the carbon cycle from satellites on a global scale.
15
Oxygen
16
Compounds The oxidation state of oxygen is −2 in almost all known compounds of oxygen. The oxidation state −1 is found in a few compounds such as peroxides.[92] Compounds containing oxygen in other oxidation states are very uncommon: −1/2 (superoxides), −1/3 (ozonides), 0 (elemental, hypofluorous acid), +1/2 (dioxygenyl), +1 (dioxygen difluoride), and +2 (oxygen difluoride).
Oxides and other inorganic compounds Water (H2O) is the oxide of hydrogen and the most familiar oxygen compound. Hydrogen atoms are covalently bonded to Water (H2O) is the most oxygen in a water molecule but also have an additional familiar oxygen compound. attraction (about 23.3 kJ·mol−1 per hydrogen atom) to an adjacent oxygen atom in a separate molecule.[93] These hydrogen bonds between water molecules hold them approximately 15% closer than what would be expected in a simple liquid with just Van der Waals forces.[94] [95] Due to its electronegativity, oxygen forms chemical bonds with almost all other elements at elevated temperatures to give corresponding oxides. However, some elements readily form oxides at standard conditions for temperature and pressure; the rusting of iron is an example. The surface of metals like aluminium and titanium are oxidized in the presence of air and become coated with a thin film of oxide that passivates the metal and slows further corrosion. Some Oxides, such as iron oxide or rust form of the transition metal oxides are found in nature as when oxygen combines with other non-stoichiometric compounds, with a slightly less elements. metal than the chemical formula would show. For example, the natural occurring FeO (wüstite) is actually written as Fe1 − xO, where x is usually around 0.05.[96] Oxygen as a compound is present in the atmosphere in trace quantities in the form of carbon dioxide (CO2). The earth's crustal rock is composed in large part of oxides of silicon (silica SiO2, found in granite and sand), aluminium (aluminium oxide Al2O3, in bauxite and corundum), iron (iron(III) oxide Fe2O3, in hematite and rust) and other metals. The rest of the Earth's crust is also made of oxygen compounds, in particular calcium carbonate (in limestone) and silicates (in feldspars). Water-soluble silicates in the form of Na4SiO4, Na2SiO3, and Na2Si2O5 are used as detergents and adhesives.[97] Oxygen also acts as a ligand for transition metals, forming metal–O2 bonds with the iridium atom in Vaska's complex,[98] with the platinum in PtF6,[99] and with the iron center of the heme group of hemoglobin.
Oxygen
17
Organic compounds and biomolecules Among the most important classes of organic compounds that contain oxygen are (where "R" is an organic group): alcohols (R-OH); ethers (R-O-R); ketones (R-CO-R); aldehydes (R-CO-H); carboxylic acids (R-COOH); esters (R-COO-R); acid anhydrides (R-CO-O-CO-R); and amides (R-C(O)-NR2). There are many important organic solvents that contain oxygen, including: acetone, methanol, ethanol, isopropanol, furan, THF, diethyl ether, dioxane, ethyl acetate, DMF, DMSO, acetic acid, and formic acid. Acetone ((CH3)2CO) and phenol (C6H5OH) are used as feeder materials in the synthesis of many different substances. Other important organic compounds that contain oxygen are: glycerol, formaldehyde, glutaraldehyde, citric acid, acetic anhydride, and acetamide. Epoxides are ethers in which the oxygen atom is part of a ring of three atoms. Oxygen
reacts
spontaneously
with
many
Acetone is an important feeder material in the chemical industry (oxygen is in red, carbon in black and hydrogen in white).
organic
compounds at or below room temperature in a process called autoxidation.[100] Most of the organic compounds that contain oxygen are not made by direct action of O2. Organic compounds important in industry and commerce that are made by direct oxidation of a precursor include ethylene oxide and peracetic acid.[97]
Oxygen represents more than 40% of the molecular mass of the ATP molecule.
The element is found in almost all biomolecules that are important to (or generated by) life. Only a few common complex biomolecules, such as squalene and the carotenes, contain no oxygen. Of the organic compounds with biological relevance, carbohydrates contain the largest proportion by mass of oxygen. All fats, fatty acids, amino acids, and proteins contain oxygen (due to the presence of carbonyl groups in these acids and their ester residues). Oxygen also occurs in phosphate (PO3−4) groups in the biologically important energy-carrying molecules ATP and ADP, in the backbone and the purines (except adenine) and pyrimidines of RNA and DNA, and in bones as calcium phosphate and hydroxylapatite.
Oxygen
18
Safety and Precautions Toxicity
Main symptoms of oxygen toxicity.
[101]
Oxygen gas (O2) can be toxic at elevated partial pressures, leading to convulsions and other health problems.[81] [102] [103] Oxygen toxicity usually begins to occur at partial pressures more than 50 kilopascals (kPa), or 2.5 times the normal sea-level O2 partial pressure of about 21 kPa. Therefore, air supplied through oxygen masks in medical applications is typically composed of 30% O2 by volume (about 30 kPa at standard pressure).[26] At one time, premature babies were placed in incubators containing O2-rich air, but this practice was discontinued after some babies were blinded by it.[26] [104]
Breathing
Oxygen toxicity occurs when lungs take in a higher than normal O2 partial pressure, which can occur in deep scuba diving.
pure
O
2
in
space
applications, such as in some modern space suits, or in early spacecraft such as Apollo, causes no damage due to the low total pressures used.[79] [105] In the case of spacesuits, the O2 partial pressure in the breathing gas is, in general, about 30 kPa (1.4 times normal), and the resulting O2 partial pressure in the astronaut's arterial blood is only marginally more than normal sea-level O2 partial pressure (see arterial blood gas).
Oxygen toxicity to the lungs and central nervous system can also occur in deep scuba diving and surface supplied diving.[26] [81] Prolonged breathing of an air mixture with an O2 partial pressure more than 60 kPa can eventually lead to permanent pulmonary fibrosis.[106] Exposure to a O2 partial pressures greater than 160 kPa may lead to convulsions (normally fatal for divers). Acute oxygen toxicity can occur by breathing an air mixture with 21% O2 at 66 m or more of depth while the same thing can occur by breathing 100% O2 at only 6 m.[106] [107] [108] [109]
Oxygen
19
Combustion and other hazards
0 0 0 OX Highly concentrated sources of oxygen promote rapid combustion. Fire and explosion hazards exist when concentrated oxidants and fuels are brought into close proximity; however, an ignition event, such as heat or a spark, is needed to trigger combustion.[110] Oxygen itself is not the fuel, but the oxidant. Combustion hazards also apply to compounds of oxygen with a high oxidative potential, such as peroxides, chlorates, nitrates, perchlorates, and dichromates because they can donate oxygen to a fire.
Pure O2 at higher than normal pressure and a spark led to a fire and the loss of the Apollo 1 crew.
Concentrated O2 will allow combustion to proceed rapidly and energetically.[110] Steel pipes and storage vessels used to store and transmit both gaseous and liquid oxygen will act as a fuel; and therefore the design and manufacture of O2 systems requires special training to ensure that ignition sources are minimized.[110] The fire that killed the Apollo 1 crew on a test launch pad spread so rapidly because the capsule was pressurized with pure O2 but at slightly more than atmospheric pressure, instead of the ⅓ normal pressure that would be used in a mission.[111] [112] Liquid oxygen spills, if allowed to soak into organic matter, such as wood, petrochemicals, and asphalt can cause these materials to detonate unpredictably on subsequent mechanical impact.[110] As with other cryogenic liquids, on contact with the human body it can cause burns to the skin and the eyes.
See also • • • • • • • • • •
Oxygen compounds Hypoxia, a lack of oxygen Hypoxia (environmental) for O2 depletion in aquatic ecology Optode for a method of measuring O2 concentration in solution Oxygen Catastrophe The sudden rise in Oxygen in the atmosphere around 2.4B years ago Oxygen isotope ratio cycle Oxygen plant Oxygen sensor Winkler test for dissolved oxygen Limiting oxygen concentration
Oxygen
References • Agostini, D.; H. Iida, and A. Takahashi (1995). "Positron emission tomography with oxygen-15 of stunned myocardium caused by coronary artery vasospasm after recovery [113] ". British Heart Journal 73 (1): 69–72. doi:10.1136/hrt.73.1.69 [114]. PMID 7888266. • Berner, Robert A. (1999-09-18). "Atmospheric oxygen over Phanerozoic time [115]". Proceedings of the National Academy of Sciences of the USA 96 (20): 10955–57. doi:10.1073/pnas.96.20.10955 [116]. PMID 10500106. http:/ / www. pnas. org/ cgi/ content/ full/ 96/ 20/ 10955. Retrieved 2007-12-16. • Britannica contributors (1911). "John Mayow [117]". Encyclopaedia Britannica (11th ed.). http:/ / www. 1911encyclopedia. org/ John_Mayow. Retrieved 2007-12-16. • Brown, Theodore L.; LeMay, Burslen (2003). Chemistry: The Central Science. Prentice Hall/Pearson Education. p. 958. ISBN 0130484504. • Cacace, Fulvio; Giulia de Petris, and Anna Troiani (2001). "Experimental Detection of Tetraoxygen". Angewandte Chemie International Edition 40 (21): 4062–65. doi:10.1002/1521-3773(20011105)40:21<4062::AID-ANIE4062>3.0.CO;2-X [118]. • Campbell, Neil A.; Reece, Jane B. (2005). Biology, 7th Edition. San Francisco: Pearson Benjamin Cummings. pp. 522–23. ISBN 0-8053-7171-0. • Chiles, James R. (2001). Inviting Disaster: Lessons from the edge of Technology: An inside look at catastrophes and why they happen. New York: HarperCollins Publishers Inc.. ISBN 0-06-662082-1. • Cook, Gerhard A.; Lauer, Carol M. (1968). "Oxygen". in Clifford A. Hampel. The Encyclopedia of the Chemical Elements. New York: Reinhold Book Corporation. pp. 499–512. LCCN 68-29938. • Crabtree, R. (2001). The Organometallic Chemistry of the Transition Metals (3rd ed.). John Wiley & Sons. pp. 152. ISBN 978-0471184232. • Daintith, John (1994). Biographical Encyclopedia of Scientists. CRC Press. ISBN 0750302879. • Desgreniers, S; Vohra, Y. K. & Ruoff, A. L. (1990). "Optical response of very high density solid oxygen to 132 GPa". J. Phys. Chem. 94: 1117–22. doi:10.1021/j100366a020 [119]. • Dole, Malcolm (1965). "The Natural History of Oxygen [120]" (PDF). The Journal of General Physiology 49: 5–27. doi:10.1085/jgp.49.1.5 [121]. PMID 5859927. http:/ / www. jgp. org/ cgi/ reprint/ 49/ 1/ 5. pdf. Retrieved 2007-12-16. • Donald, Kenneth (1992). Oxygen and the Diver. England: SPA in conjunction with K. Donald. ISBN 1854211765. • Emsley, John (2001). "Oxygen". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. pp. 297–304. ISBN 0198503407. • Evans, David Hudson; Claiborne, James B. (2006). The Physiology of Fishes. CRC Press. pp. 88. ISBN 0849320224. • Fenical, William (September 1983). "Marine Plants: A Unique and Unexplored Resource [122] ". Plants: the potentials for extracting protein, medicines, and other useful chemicals (workshop proceedings). DIANE Publishing. p. 147. ISBN 1428923977. http:/ / books. google. com/ books?id=g6RfkqCUQyQC& pg=PA147& dq=oxygen+ percent+ algae+ plants& sig=4tJv81njIlr7qsWD95pHcuRlffc#PPA147,M1.
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• Freeman, Scott (2005). Biological Science, 2nd. Upper Saddle River, NJ: Pearson Prentice Hall. pp. 214, 586. ISBN 0-13-140941-7. • Greenwood, Norman N.; Earnshaw, A. (1997), Chemistry of the Elements (2nd ed.), Oxford: Butterworth-Heinemann, ISBN 0-7506-3365-4 • Harrison, Roy M. (1990). Pollution: Causes, Effects & Control (2nd ed.). Cambridge: Royal Society of Chemistry. ISBN 0-85186-283-7. • Hirayama, Osamu; Kyoko Nakamura, Syoko Hamada and Yoko Kobayasi (1994-02-). "Singlet oxygen quenching ability of naturally occurring carotenoids [123]". Lipids (Springer Berlin / Heidelberg) 29 (2): 149–50. doi:10.1007/BF02537155 [124]. http:/ / www. springerlink. com/ content/ d67361221v12082m/ . Retrieved 2007-12-15. • How Products are Made contributors (2002). "Oxygen [125]". How Products are Made. The Gale Group, Inc. http:/ / www. answers. com/ topic/ oxygen. Retrieved 2007-12-16. [126]
• Jastrow, Joseph (1936). Story of Human Error . Ayer Publishing. pp. 171. ISBN 0836905687. http:/ / books. google. com/ books?id=tRUO45YfCHwC& pg=PA171& lpg=PA171& dq=philo+ of+ byzantium+ combustion& source=web& ots=Nv2brEX543& sig=jBvqi2t4sg5S0RUEX864xIgdfCE#PPA171,M1. Retrieved 2007-12-16. • Krieger-Liszkay, Anja (2005). "Singlet oxygen production in photosynthesis [127]". Journal of Experimental Botanics (Oxford Journals) 56: 337–46. doi:10.1093/jxb/erh237 [128]. PMID 15310815. http:/ / jxb. oxfordjournals. org/ cgi/ content/ full/ 56/ 411/ 337. Retrieved 2007-12-16. • Lide, David R. (2003). "Section 4, Properties of the Elements and Inorganic Compounds; Melting, boiling, and critical temperatures of the elements". CRC Handbook of Chemistry and Physics (84th ed.). Boca Raton, Florida: CRC Press. • Lundegaard, Lars F.; Weck, Gunnar; McMahon, Malcolm I.; Desgreniers, Serge and Loubeyre, Paul (2006). "Observation of an O8 molecular lattice in the phase of solid oxygen [129]". Nature 443: 201–04. doi:10.1038/nature05174 [130]. http:/ / www. nature. com/ nature/ journal/ v443/ n7108/ abs/ nature05174. html. Retrieved 2008-01-10. • Maksyutenko, P.; T. R. Rizzo, and O. V. Boyarkin (2006). "A direct measurement of the dissociation energy of water". J. Chem. Phys. 443: 125. doi:10.1063/1.2387163 [131]. • Meyer, B.S. (September 19–21, 2005). "Nucleosynthesis and Galactic Chemical Evolution of the Isotopes of Oxygen [132]" (PDF). Workgroup on Oxygen in the Earliest Solar System [133] . Proceedings of the NASA Cosmochemistry Program and the Lunar and Planetary Institute. Gatlinburg, Tennessee. 9022. http:/ / www. lpi. usra. edu/ meetings/ ess2005/ pdf/ 9022. pdf. Retrieved 2007-01-22. • Miller, J.R.; Berger, M.; Alonso, L.; Cerovic, Z.; Goulas, Y.; Jacquemoud, S.; Louis, J.; Mohammed, G.; Moya, I.; Pedros, R.; Moreno, J.F.; Verhoef, W.; Zarco-Tejada, P.J.. "Progress on the development of an integrated canopy fluorescence model [134]". Geoscience and Remote Sensing Symposium, 2003. IGARSS '03. Proceedings. 2003 IEEE International. • Morris, Richard (2003) (Hardback). The last sorcerers: The path from alchemy to the periodic table. Washington, D.C.: Joseph Henry Press. ISBN 0309089050. • Parks, G. D.; Mellor, J. W. (1939). Mellor's Modern Inorganic Chemistry (6th ed.). London: Longmans, Green and Co. • Priestley, Joseph (1775). "An Account of Further Discoveries in Air [135]". Philosophical Transactions 65: 384–94. doi:10.1098/rstl.1775.0039 [136]. http:/ / links. jstor. org/
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sici?sici=0260-7085%281775%2965%3C384%3AAAOFDI%3E2. 0. CO%3B2-N. Retrieved 2007-12-16. • Raven, Peter H.; Ray F. Evert, Susan E. Eichhorn (2005). Biology of Plants, 7th Edition. New York: W.H. Freeman and Company Publishers. pp. 115–27. ISBN 0-7167-1007-2. • Roscoe, Henry Enfield; Schorlemmer, Carl (1883). A Treatise on Chemistry. D. Appleton and Co.. pp. 38. • Shimizu, K.; Suhara, K., Ikumo, M., Eremets, M. I. & Amaya, K. (1998). "Superconductivity in oxygen". Nature 393: 767–69. doi:10.1038/31656
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• Smart, Lesley E.; Elaine A. Moore (2005). Solid State Chemistry: An Introduction (3rd ed.). CRC Press. pp. 214. ISBN 978-0748775163. • Stwertka, Albert (1998). Guide to the Elements (Revised ed.). Oxford University Press. ISBN 0-19-508083-1. • Walker, J. (1980). "The oxygen cycle". in Hutzinger O.. Handbook of Environmental Chemistry. Volume 1. Part A: The natural environment and the biogeochemical cycles. Berlin; Heidelberg; New York: Springer-Verlag. pp. 258. ISBN 0387096884. • Wentworth Jr., Paul; Jonathan E. McDunn, Anita D. Wentworth, Cindy Takeuchi, Jorge Nieva, Teresa Jones, Cristina Bautista, Julie M. Ruedi, Abel Gutierrez, Kim D. Janda, Bernard M. Babior, Albert Eschenmoser, Richard A. Lerner (2002-12-13). "Evidence for Antibody-Catalyzed Ozone Formation in Bacterial Killing and Inflammation". Science 298 (5601): 2195–219. doi:10.1126/science.1077642 [138]. PMID 12434011. • Werley, Barry L. (Edtr.) (1991). "Fire Hazards in Oxygen Systems". ASTM Technical Professional training. Philadelphia: ASTM International Subcommittee G-4.05. • World of Chemistry contributors (2005). "John Mayow [139]". World of Chemistry. Thomson Gale. http:/ / www. bookrags. com/ John_Mayow. Retrieved 2007-12-16.
External links • Oxidizing Agents > Oxygen [140] • Oxygen (O2) Properties, Uses, Applications • Roald Hoffmann article on "The Story of O" • WebElements.com – Oxygen [143]
[141] [142]
pnb:نجیسکآ
References [1] Emsley 2001, p.297
" Oxygen (http:/ / periodic. lanl. gov/ elements/ 8. html)". Los Alamos National Laboratory. . Retrieved 2007-12-16. [3] Cook & Lauer 1968, p.500 [4] NASA (2007-09-27). " NASA Research Indicates Oxygen on Earth 2.5 Billion Years Ago (http:/ / www. nasa. gov/ home/ hqnews/ 2007/ sep/ HQ_07215_Timeline_of_Oxygen_on_Earth. html)". Press release. . Retrieved 2008-03-13. [5] " Atomic oxygen erosion (http:/ / www. spenvis. oma. be/ spenvis/ help/ background/ atmosphere/ erosion. html)". . Retrieved 2009-08-08. [6] Mellor 1939 [7] " Molecular Orbital Theory (http:/ / chemed. chem. purdue. edu/ genchem/ topicreview/ bp/ ch8/ mo. html#bond)". Purdue University. . Retrieved 2008-01-28.
Oxygen [8] Pauling, L. (1960). The nature of the chemical bond and the structure of molecules and crystals : an introduction to modern structural chemistry (3rd ed.). Ithaca, N.Y.: Cornell University Press. [9] Jakubowski, Henry. " Biochemistry Online (http:/ / employees. csbsju. edu/ hjakubowski/ classes/ ch331/ bcintro/ default. html)". Saint John's University. . Retrieved 2008-01-28. [10] An orbital is a concept from quantum mechanics that models an electron as a wave-like particle that has a spacial distribution about an atom or molecule. [11] Emsley 2001, p.303 [12] " Demonstration of a bridge of liquid oxygen supported against its own weight between the poles of a powerful magnet (http:/ / genchem. chem. wisc. edu/ demonstrations/ Gen_Chem_Pages/ 0809bondingpage/ liquid_oxygen. htm)". University of Wisconsin-Madison Chemistry Department Demonstration lab. . Retrieved 2007-12-15. [13] Oxygen's paramagnetism can be used analytically in paramagnetic oxygen gas analysers that determine the purity of gaseous oxygen. (" Company literature of Oxygen analyzers (triplet) (http:/ / www. servomex. com/ oxygen_gas_analyser. html)". Servomex. . Retrieved 2007-12-15.) [14] Krieger-Liszkay 2005, 337-46 [15] Harrison 1990 [16] Wentworth 2002 [17] Hirayama 1994, 149-150 [18] Chieh, Chung. " Bond Lengths and Energies (http:/ / www. science. uwaterloo. ca/ ~cchieh/ cact/ c120/ bondel. html)". University of Waterloo. . Retrieved 2007-12-16. [19] Stwertka 1998, p.48 [20] Stwertka 1998, p.49 [21] Cacace 2001, 4062 [22] Ball, Phillip (2001-09-16). " New form of oxygen found (http:/ / www. nature. com/ news/ 2001/ 011122/ pf/ 011122-3_pf. html)". Nature News. . Retrieved 2008-01-09. [23] Lundegaard 2006, 201–04 [24] Desgreniers 1990, 1117–22 [25] Shimizu 1998, 767–69 [26] Emsley 2001, p.299 [27] " Air solubility in water (http:/ / www. engineeringtoolbox. com/ air-solubility-water-d_639. html)". The Engineering Toolbox. . Retrieved 2007-12-21. [28] Evans & Claiborne 2006, 88 [29] Lide 2003, Section 4 [30] " Overview of Cryogenic Air Separation and Liquefier Systems (http:/ / www. uigi. com/ cryodist. html)". Universal Industrial Gases, Inc.. . Retrieved 2007-12-15. [31] " Liquid Oxygen Material Safety Data Sheet (https:/ / www. mathesontrigas. com/ pdfs/ msds/ 00225011. pdf)" (PDF). Matheson Tri Gas. . Retrieved 2007-12-15. [32] " Oxygen Nuclides / Isotopes (http:/ / environmentalchemistry. com/ yogi/ periodic/ O-pg2. html)". EnvironmentalChemistry.com. . Retrieved 2007-12-17. [33] Meyer 2005, 9022 [34] " NUDAT 13O (http:/ / www. nndc. bnl. gov/ nudat2/ decaysearchdirect. jsp?nuc=13O& unc=nds)". . Retrieved 2009-07-06. [35] " NUDAT 14O (http:/ / www. nndc. bnl. gov/ nudat2/ decaysearchdirect. jsp?nuc=14O& unc=nds)". . Retrieved 2009-07-06. [36] " NUDAT 15O (http:/ / www. nndc. bnl. gov/ nudat2/ decaysearchdirect. jsp?nuc=15O& unc=nds)". . Retrieved 2009-07-06. [37] Emsley 2001, p.298 [38] Figures given are for values up to 50 miles (80 km) above the surface [39] From The Chemistry and Fertility of Sea Waters by H.W. Harvey, 1955, citing C.J.J. Fox, "On the coefficients of absorption of atmospheric gases in sea water", Publ. Circ. Cons. Explor. Mer, no. 41, 1907. Harvey however notes that according to later articles in Nature the values appear to be about 3% too high. [40] Emsley 2001, p.301 [41] Fenical 1983, "Marine Plants" [42] Brown 2003, 958 [43] Thylakoid membranes are part of chloroplasts in algae and plants while they simply are one of many membrane structures in cyanobacteria. In fact, chloroplasts are thought to have evolved from cyanobacteria that were once symbiotic partners with the progenerators of plants and algae. [44] Raven 2005, 115–27
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Oxygen [45] Water oxidation is catalyzed by a manganese-containing enzyme complex known as the oxygen evolving complex (OEC) or water-splitting complex found associated with the lumenal side of thylakoid membranes. Manganese is an important cofactor, and calcium and chloride are also required for the reaction to occur.(Raven 2005) [46] CO2 is released from another part of hemoglobin (see Bohr effect) [47] "For humans, the normal volume is 6-8 liters per minute." (http:/ / www. patentstorm. us/ patents/ 6224560-description. html) [48] (1.8 grams/min/person)×(60 min/h)×(24 h/day)×(365 days/year)×(6.6 billion people)/1,000,000 g/t=6.24 billion tonnes [49] Campbell 2005, 522–23 [50] Freeman 2005, 214, 586 [51] Berner 1999, 10955–57 [52] Dole 1965, 5–27 [53] Jastrow 1936, 171 [54] Cook & Lauer 1968, p.499. [55] Britannica contributors 1911, "John Mayow" [56] World of Chemistry contributors 2005, "John Mayow" [57] Morris 2003 [58] Emsley 2001, p.300 [59] Priestley 1775, 384–94 [60] DeTurck, Dennis; Gladney, Larry and Pietrovito, Anthony (1997). " The Interactive Textbook of PFP96 (http:/ / www. physics. upenn. edu/ courses/ gladney/ mathphys/ Contents. html)". University of Pennsylvania. . Retrieved 2008-01-28. [61] Roscoe 1883, 38 [62] However, these results were mostly ignored until 1860. Part of this rejection was due to the belief that atoms of one element would have no chemical affinity towards atoms of the same element, and part was due to apparent exceptions to Avogadro's law that were not explained until later in terms of dissociating molecules. [63] Daintith 1994, p.707 [64] Poland - Culture, Science and Media. Condensation of oxygen and nitrogen (http:/ / www. poland. gov. pl/ Karol,Olszewski,and,Zygmunt,Wroblewski:,condensation,of,oxygen,and,nitrogen,1987. html). Retrieved on 2008-10-04. [65] How Products are Made contributors, "Oxygen" [66] " Goddard-1926 (http:/ / grin. hq. nasa. gov/ ABSTRACTS/ GPN-2002-000132. html)". NASA. . Retrieved 2007-11-18. [67] " Non-Cryogenic Air Separation Processes (http:/ / www. uigi. com/ noncryo. html)". UIG Inc.. 2003. . Retrieved 2007-12-16. [68] Space Shuttle Use of Propellants and Fluids (http:/ / www-pao. ksc. nasa. gov/ kscpao/ nasafact/ ps/ SSP. ps), National Aeronautics and Space Administration, 2001-09, , retrieved 2007-12-16, "NASAFacts FS-2001-09-015-KSC" [69] Cook & Lauer 1968, p.510 [70] Sim MA, Dean P, Kinsella J, Black R, Carter R, Hughes M (2008). "Performance of oxygen delivery devices when the breathing pattern of respiratory failure is simulated". Anaesthesia 63 (9): 938–40. doi: 10.1111/j.1365-2044.2008.05536.x (http:/ / dx. doi. org/ 10. 1111/ j. 1365-2044. 2008. 05536. x). PMID 18540928. [71] Stephenson RN, Mackenzie I, Watt SJ, Ross JA (1996). " Measurement of oxygen concentration in delivery systems used for hyperbaric oxygen therapy (http:/ / archive. rubicon-foundation. org/ 2245)". Undersea Hyperb Med 23 (3): 185–8. PMID 8931286. . Retrieved 2008-09-22. [72] Undersea and Hyperbaric Medical Society. " Indications for hyperbaric oxygen therapy (http:/ / www. uhms. org/ Default. aspx?tabid=270)". . Retrieved 2008-09-22. [73] Undersea and Hyperbaric Medical Society. " Carbon Monoxide (http:/ / www. uhms. org/ ResourceLibrary/ Indications/ CarbonMonoxidePoisoning/ tabid/ 272/ Default. aspx)". . Retrieved 2008-09-22. [74] Piantadosi CA (2004). " Carbon monoxide poisoning (http:/ / archive. rubicon-foundation. org/ 4002)". Undersea Hyperb Med 31 (1): 167–77. PMID 15233173. . Retrieved 2008-09-22. [75] Hart GB, Strauss MB (1990). " Gas Gangrene - Clostridial Myonecrosis: A Review (http:/ / archive. rubicon-foundation. org/ 4428)". J. Hyperbaric Med 5 (2): 125–144. . Retrieved 2008-09-22. [76] Zamboni WA, Riseman JA, Kucan JO (1990). " Management of Fournier's Gangrene and the role of Hyperbaric Oxygen (http:/ / archive. rubicon-foundation. org/ 4431)". J. Hyperbaric Med 5 (3): 177–186. . Retrieved 2008-09-22.
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Oxygen [77] Undersea and Hyperbaric Medical Society. " Decompression Sickness or Illness and Arterial Gas Embolism (http:/ / www. uhms. org/ ResourceLibrary/ Indications/ DecompressionSickness/ tabid/ 275/ Default. aspx)". . Retrieved 2008-09-22. [78] Acott, C. (1999). " A brief history of diving and decompression illness (http:/ / archive. rubicon-foundation. org/ 6004)". South Pacific Underwater Medicine Society journal 29 (2). ISSN 0813-1988 (http:/ / worldcat. org/ issn/ 0813-1988). OCLC 16986801 (http:/ / worldcat. org/ oclc/ 16986801). . Retrieved 2008-09-22. [79] Morgenthaler GW, Fester DA, Cooley CG (1994). "As assessment of habitat pressure, oxygen fraction, and EVA suit design for space operations". Acta Astronaut 32 (1): 39–49. doi: 10.1016/0094-5765(94)90146-5 (http:/ / dx. doi. org/ 10. 1016/ 0094-5765(94)90146-5). PMID 11541018. [80] Webb JT, Olson RM, Krutz RW, Dixon G, Barnicott PT (1989). "Human tolerance to 100% oxygen at 9.5 psia during five daily simulated 8-hour EVA exposures". Aviat Space Environ Med 60 (5): 415–21. PMID 2730484. [81] Acott, C. (1999). " Oxygen toxicity: A brief history of oxygen in diving (http:/ / archive. rubicon-foundation. org/ 6014)". South Pacific Underwater Medicine Society journal 29 (3). ISSN 0813-1988 (http:/ / worldcat. org/ issn/ 0813-1988). OCLC 16986801 (http:/ / worldcat. org/ oclc/ 16986801). . Retrieved 2008-09-21. [82] Longphre, J. M.; P. J. DeNoble; R. E. Moon; R. D. Vann; J. J. Freiberger (2007). " First aid normobaric oxygen for the treatment of recreational diving injuries. (http:/ / archive. rubicon-foundation. org/ 5514)". Undersea Hyperb Med. 34 (1): 43–49. ISSN 1066-2936 (http:/ / worldcat. org/ issn/ 1066-2936). OCLC 26915585 (http:/ / worldcat. org/ oclc/ 26915585). PMID 17393938. . Retrieved 2008-09-21. [83] The reason is that increasing the proportion of oxygen in the breathing gas at low pressure acts to augment the inspired O2 partial pressure nearer to that found at sea-level. [84] " NTSB Summary report (http:/ / www. ntsb. gov/ NTSB/ brief. asp?ev_id=20001208X05743& key=1)". National Transportation Safety Board. . Retrieved 2007-12-16.) [85] Bren, Linda (November–December 2002). " Oxygen Bars: Is a Breath of Fresh Air Worth It? (http:/ / www. fda. gov/ Fdac/ features/ 2002/ 602_air. html)". FDA Consumer magazine. U.S. Food and Drug Administration. . Retrieved 2007-12-23. [86] " Ergogenic Aids (http:/ / www. pponline. co. uk/ encyc/ 1008. htm)". Peak Performance Online. . Retrieved 2008-01-04. [87] " George Goble's extended home page (mirror) (http:/ / www. bkinzel. de/ misc/ ghg/ index. html)". . [88] Cook & Lauer 1968, p.508 [89] Emsley 2001, p.304 [90] Hand, Eric (2008-03-13). " The Solar System's first breath (http:/ / www. nature. com/ news/ 2008/ 080313/ full/ 452259a. html)". Nature 452: 259. doi: 10.1038/452259a (http:/ / dx. doi. org/ 10. 1038/ 452259a). . Retrieved 2009-03-18. [91] Miller et al. 2003 [92] Greenwood & Earnshaw 1997, 28 [93] Maksyutenko et al. 2006 [94] Chaplin, Martin (2008-01-04). " Water Hydrogen Bonding (http:/ / www. lsbu. ac. uk/ water/ hbond. html)". . Retrieved 2008-01-06. [95] Also, since oxygen has a higher electronegativity than hydrogen, the charge difference makes it a polar molecule. The interactions between the different dipoles of each molecule cause a net attraction force. [96] Smart 2005, 214 [97] Cook & Lauer 1968, p.507 [98] Crabtree 2001, 152 [99] Cook & Lauer 1968, p.505 [100] Cook & Lauer 1968, p.506 [101] Dharmeshkumar N Patel, Ashish Goel, SB Agarwal, Praveenkumar Garg, Krishna K Lakhani (2003). " Oxygen Toxicity (http:/ / medind. nic. in/ jac/ t03/ i3/ jact03i3p234. pdf)". Indian Academy of Clinical Medicine 4 (3): 234. . [102] Since O2's partial pressure is the fraction of O2 times the total pressure, elevated partial pressures can occur either from high O2 fraction in breathing gas or from high breathing gas pressure, or a combination of both. [103] Cook & Lauer 1968, p.511 [104] Drack AV (1998). "Preventing blindness in premature infants". N. Engl. J. Med. 338 (22): 1620–1. doi: 10.1056/NEJM199805283382210 (http:/ / dx. doi. org/ 10. 1056/ NEJM199805283382210). PMID 9603802. [105] Wade, Mark (2007). " Space Suits (http:/ / www. astronautix. com/ craftfam/ spasuits. htm)". Encyclopedia Astronautica. . Retrieved 2007-12-16. [106] Wilmshurst P (1998). " Diving and oxygen (http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?tool=pmcentrez& artid=1114047)". BMJ 317 (7164): 996–9. PMID 9765173. [107] Donald 1992
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[108] Donald K. W. (1947). " Oxygen Poisoning in Man: Part I (http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?tool=pmcentrez& artid=2053251)". Br Med J 1 (4506): 667–72. [109] Donald K. W. (1947). " Oxygen Poisoning in Man: Part II (http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?tool=pmcentrez& artid=2053400)". Br Med J 1 (4507): 712–7. [110] Werley 1991 [111] No single ignition source of the fire was conclusively identified, although some evidence points to arc from an electrical spark). (Report of Apollo 204 Review Board NASA Historical Reference Collection, NASA History Office, NASA HQ, Washington, DC) [112] Chiles 2001 [113] http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?tool=pmcentrez& artid=483759 [114] http:/ / dx. doi. org/ 10. 1136%2Fhrt. 73. 1. 69 [115] http:/ / www. pnas. org/ cgi/ content/ full/ 96/ 20/ 10955 [116] http:/ / dx. doi. org/ 10. 1073%2Fpnas. 96. 20. 10955 [117] http:/ / www. 1911encyclopedia. org/ John_Mayow [118] http:/ / dx. doi. org/ 10. 1002%2F1521-3773%2820011105%2940%3A21%3C4062%3A%3AAID-ANIE4062%3E3. 0. CO%3B2-X [119] http:/ / dx. doi. org/ 10. 1021%2Fj100366a020 [120] http:/ / www. jgp. org/ cgi/ reprint/ 49/ 1/ 5. pdf [121] http:/ / dx. doi. org/ 10. 1085%2Fjgp. 49. 1. 5 [122] http:/ / books. google. com/ books?id=g6RfkqCUQyQC& pg=PA147& dq=oxygen+ percent+ algae+ plants& sig=4tJv81njIlr7qsWD95pHcuRlffc#PPA147,M1 [123] http:/ / www. springerlink. com/ content/ d67361221v12082m/ [124] http:/ / dx. doi. org/ 10. 1007%2FBF02537155 [125] http:/ / www. answers. com/ topic/ oxygen [126] http:/ / books. google. com/ books?id=tRUO45YfCHwC& pg=PA171& lpg=PA171& dq=philo+ of+ byzantium+ combustion& source=web& ots=Nv2brEX543& sig=jBvqi2t4sg5S0RUEX864xIgdfCE#PPA171,M1 [127] [128] [129] [130] [131] [132] [133] [134] [135] [136] [137] [138] [139] [140] [141] [142] [143]
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Article Sources and Contributors
Article Sources and Contributors Oxygen Source: http://en.wikipedia.org/w/index.php?oldid=308017614 Contributors: 0612, 123qwe, 1266asdsdjapg, 1297, Abarenbo, Acalamari, Acroterion, Acs4b, Adambiswanger1, Adashiel, Addshore, Adrian, AdultSwim, Aeros320, AgainErick, Ahoerstemeier, AidepikiW kcuF, Aitias, Alex.muller, AlexG, Alexeymorgunov, Alexf, AlexiusHoratius, Algont, Alias Flood, Alison, All Is One, Alsandro, Amphetamine Analogue, Ancheta Wis, AndonicO, Andre Engels, Andres, Andrewlp1991, Andros 1337, AngelOfSadness, AngryParsley, Angusmclellan, Ann Stouter, Anna512, AnonMoos, Antandrus, Anthony Appleyard, Arcadian, Archimerged, Ardric47, Arjun01, Arkuat, ArnoldReinhold, Arnon Chaffin, Atemperman, AuburnPilot, Axlq, AzaToth, BANZ111, BHS Sux, Bachrach44, Badocter, BalazsH, Ballsonyourwalls, Balthazarduju, BanyanTree, Bart133, Bbatsell, Bbi5291, Bboy14, Beetstra, Beland, Benbest, Benjah-bmm27, Benjiboi, Bensderbest, Bevo, Bggoldie, Bhadani, BigCow, Bigbuck, Billcurtis, Blackjack3, BlueEarth, BlueMoonlet, Bluezy, Bobak, Bobo192, Bomac, Bongwarrior, Bonjour amis, Bornfury, Borovy3488, Bradkittenbrink, Bryan Derksen, Buchanan-Hermit, Buckthebronco, Burntsauce, Burzmali, Buzzgrav08, C777, CJLL Wright, CWii, CYD, Cactus.man, Caesura, Caltas, CambridgeBayWeather, Can't sleep, clown will eat me, CanadianCaesar, Candlewicke, CanisRufus, Carlo.milanesi, Carloseduardo, Carnildo, Casliber, Catslash, CattleGirl, Causesobad, Cd12holden, Cdc, CelticJobber, Cfw Master, Chameleon, Charleythegodfather, Chcknpie04, Chilisauce2727, Chlämens, Cholmes75, Chowbok, Chris Dybala, Chris the speller, Chrislk02, Christian List, Christopherlin, Ck lostsword, Clivegrey, Cmapm, Cnaude, Colbuckshot, ColdFeet, Cometstyles, CommonsDelinker, Conn, Kit, Conversion script, Corpx, Cosmium, Cryptic C62, Crystallina, Ctbolt, Curps, D, DJ Clayworth, DRosenbach, DVD R W, Dac107, Damieng, Dan56, Dana boomer, DancingPenguin, Daniel Case, DanielCD, Danny, Dantheman531, Danyg, Dark Mage, Darrien, Dauno, David Latapie, Davidj1991, Davumaya, Dawn Bard, Ddday-z, DeadEyeArrow, Deglr6328, Deli nk, Delta G, Demoscn, Deonfjw, DerHexer, Derek Ross, Derek.cashman, Deryck Chan, Devl2666, Digitalme, Dillard421, Dirac66, DirectEdge, Discospinster, Dmz5, Doct.proloy, Dolive21, DomCleal, Dominus, DonSiano, Donarreiskoffer, Dragonmaster84, Dravick, Dreadstar, Drini, Droll, Drphilharmonic, Dsyzdek, Dwmyers, Dycedarg, Dysepsion, EEMIV, EL Willy, Ecophreek, EdC, Eddideigel, Edgar181, Edsanville, Edward, Egil, Eilatybartfast, El C, Eldin raigmore, Eleassar, Eliz81, Elkman, Emperorbma, Enemyunknown, Eng02019, Enigmaman, Eric Forste, Eric Kvaalen, Eric119, Erik Zachte, Etaoin, Ethel Aardvark, Euyyn, Evercat, Everyking, Evil Monkey, Ewen, Ex nihil, Exarion, Eye.earth, FTGHSmith, Fabartus, Faithlessthewonderboy, Femto, Finell, FisherQueen, FrancoGG, Francs2000, Freakofnurture, Fredrik, Frencheigh, FreplySpang, Frymaster, Funky Monkey, Fvw, G. 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Image Sources, Licenses and Contributors file:cubic.svg Source: http://en.wikipedia.org/w/index.php?title=File:Cubic.svg License: GNU Free Documentation License Contributors: User:Stannered file:Electron shell 008 Oxygen.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_008_Oxygen.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 file:Liquid Oxygen2.gif Source: http://en.wikipedia.org/w/index.php?title=File:Liquid_Oxygen2.gif License: unknown Contributors: User:Materialscientist Image:Ozone-montage.png Source: http://en.wikipedia.org/w/index.php?title=File:Ozone-montage.png License: Public Domain Contributors: Benjah-bmm27, Teebeutel Image:Evolved star fusion shells.svg Source: http://en.wikipedia.org/w/index.php?title=File:Evolved_star_fusion_shells.svg License: unknown Contributors: User:Rursus File:WOA05 sea-surf O2 AYool.png Source: http://en.wikipedia.org/w/index.php?title=File:WOA05_sea-surf_O2_AYool.png License: unknown Contributors: User:Plumbago Image:Simple photosynthesis overview.svg Source: http://en.wikipedia.org/w/index.php?title=File:Simple_photosynthesis_overview.svg License: unknown Contributors: User:Maveric149, User:Yerpo
27
Image Sources, Licenses and Contributors File:Auto-and heterotrophs.png Source: http://en.wikipedia.org/w/index.php?title=File:Auto-and_heterotrophs.png License: unknown Contributors: User:Mikael Häggström Image:Oxygenation-atm.svg Source: http://en.wikipedia.org/w/index.php?title=File:Oxygenation-atm.svg License: unknown Contributors: Heinrich D. Holland Image:Philos experiment of the burning candle.PNG Source: http://en.wikipedia.org/w/index.php?title=File:Philos_experiment_of_the_burning_candle.PNG License: Public Domain Contributors: Wilhelm Schmidt Image:Georg Ernst Stahl.png Source: http://en.wikipedia.org/w/index.php?title=File:Georg_Ernst_Stahl.png License: Public Domain Contributors: Polarlys Image:Carl Wilhelm Scheele from Familj-Journalen1874.png Source: http://en.wikipedia.org/w/index.php?title=File:Carl_Wilhelm_Scheele_from_Familj-Journalen1874.png License: Public Domain Contributors: Celsius, Crux, Den fjättrade ankan, Sanao, 1 anonymous edits Image:PriestleyFuseli.jpg Source: http://en.wikipedia.org/w/index.php?title=File:PriestleyFuseli.jpg License: unknown Contributors: Turner, Charles , 1774 - 1857 (Engraver); Fuseli, Henry, 1741 - 1825 (Painter) Image:Antoine lavoisier.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Antoine_lavoisier.jpg License: Public Domain Contributors: Kilom691, Matanya, Siebrand Image:Goddard and Rocket.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Goddard_and_Rocket.jpg License: Public Domain Contributors: Unknown Image:209px-Hofmann voltameter.fr.version.svg.png Source: http://en.wikipedia.org/w/index.php?title=File:209px-Hofmann_voltameter.fr.version.svg.png License: GNU Free Documentation License Contributors: Itub, Lokal Profil, Mion Image:Home oxygen concentrator.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Home_oxygen_concentrator.jpg License: Creative Commons Attribution-Sharealike 2.0 Contributors: Original uploader was GiollaUidir at en.wikipedia Image:Wisoff on the Arm - GPN-2000-001069.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Wisoff_on_the_Arm_-_GPN-2000-001069.jpg License: Public Domain Contributors: NASA Image:Clabecq JPG01.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Clabecq_JPG01.jpg License: unknown Contributors: user: Jean-Pol GRANDMONT Image:Phanerozoic Climate Change.png Source: http://en.wikipedia.org/w/index.php?title=File:Phanerozoic_Climate_Change.png License: unknown Contributors: Royer, Dana L., Robert A. Berner, Isabel P. Montañez, Neil J. Tabor, and David J. Beerling Image:Stilles Mineralwasser.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Stilles_Mineralwasser.jpg License: GNU Free Documentation License Contributors: Walter J. Pilsak, Waldsassen, Germany Image:Rust screw.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Rust_screw.jpg License: Creative Commons Attribution 2.0 Contributors: User:Paulnasca. Original uploader was Paulnasca at en.wikipedia Image:Acetone-3D-vdW.png Source: http://en.wikipedia.org/w/index.php?title=File:Acetone-3D-vdW.png License: Public Domain Contributors: Ben Mills Image:ATP structure.svg Source: http://en.wikipedia.org/w/index.php?title=File:ATP_structure.svg License: Public Domain Contributors: w:User:MysidUser:Mysid File:Symptoms of oxygen toxicity.png Source: http://en.wikipedia.org/w/index.php?title=File:Symptoms_of_oxygen_toxicity.png License: Public Domain Contributors: Mikael Häggström Image:Scuba-diving.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Scuba-diving.jpg License: Creative Commons Attribution-Sharealike 2.5 Contributors: BLueFiSH.as, Civertan, Fschoenm, Man vyi, Wikipeder, 3 anonymous edits Image:NFPA 704.svg Source: http://en.wikipedia.org/w/index.php?title=File:NFPA_704.svg License: Public Domain Contributors: User:Denelson83 Image:Apollo 1 fire.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Apollo_1_fire.jpg License: Public Domain Contributors: NASA
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28
Fluorine
1
Fluorine Oxygen ← Fluorine → Neon ↑ F ↓ Cl
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
9F Periodic table
Appearance Yellowish brown gas General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensityMelting pointBoiling pointCritical pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
38
44
50
58
69
85
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 3374.2 kJ·mol−1 3rd: 6050.4 kJ·mol−1Covalent radiusVan der Waals radius Miscellaneous Crystal structureMagnetic orderingThermal conductivityCAS registry number Most stable isotopes Main article: Isotopes of Fluorine iso 18
F
19
F
N.A. syn
100%
half-life 109.77 min
19
DM
DE (MeV)
DP
β+ (97%)
0.64
18
ε (3%)
1.656
18
F is stable with 10 neutron
O O
Fluorine Fluorine, F, 9 halogen17, 2, p18.9984032(5) g·mol−1 1s2 2s2 2p5 2, 7 (Image) gas (0 °C, 101.325 kPa) 1.7 g/L 53.53 K,−219.62 °C,−363.32 °F 85.03 K,−188.12 °C,−306.62 °F 144.13 K, 5.172 MPa (F2) 0.510 kJ·mol−1 (F2) 6.62 kJ·mol−1 (25 °C) (F2) −1 −1 31.304 J·mol ·K −1 (Weaklyacidic oxide) 3.98 (Pauling scale) 1st: 1681.0 kJ·mol−157±3 pm −1 −1 (see covalent radius of fluorine)147 pm cubic nonmagnetic (300 K) 27.7 m W·m ·K 7782-41-4 Fluorine is the chemical element with atomic number 9, represented by the symbol F. Fluorine forms a single bond with itself in elemental form, resulting in the diatomic F2 molecule. F2 is a supremely reactive, poisonous, pale, yellowish brown gas. Elemental fluorine is the most chemically reactive and electronegative of all the elements. For example, it will readily "burn" hydrocarbons at room temperature, in contrast to the combustion of hydrocarbons by oxygen, which requires an input of energy with a spark. Therefore, molecular fluorine is highly dangerous, more so than other halogens such as the poisonous chlorine gas. Fluorine's highest electronegativity and small atomic radius give unique properties to many of its compounds. For example, the enrichment of 235U, the principal nuclear fuel, relies on the volatility of UF6. Also, the carbon–fluorine bond is one of the strongest bonds in organic chemistry. This contributes to the stability and persistence of fluoroalkane based organofluorine compounds, such as PTFE/(Teflon) and PFOS. The carbon–fluorine bond's inductive effects result in the strength of many fluorinated acids, such as triflic acid and trifluoroacetic acid. Drugs are often fluorinated at biologically reactive positions, to prevent their metabolism and prolong their half-lives.
Characteristics F2 is a corrosive pale yellow or brown[1] gas that is a powerful oxidizing agent. It is the most reactive and most electronegative of all the elements on the classic Pauling scale (4.0), and readily forms compounds with most other elements. It has an oxidation number -1, except when bonded to another fluorine in F2 which gives it an oxidation number of 0. Fluorine even combines with the noble gases argon, krypton, xenon, and radon. Even in dark, cool conditions, fluorine reacts explosively with hydrogen. The reaction with hydrogen occurs even at extremely low temperatures, using liquid hydrogen and solid fluorine. It is so reactive that metals, and even water, as well as other substances, burn with a bright flame in a jet of fluorine gas. In moist air it reacts with water to form the also dangerous hydrofluoric acid. Fluorides are compounds that combine fluorine with some positively charged counterpart. They often consist of crystalline ionic salts. Fluorine compounds with metals are among the most stable of salts. Hydrogen fluoride is a weak acid when dissolved in water, but is still very corrosive and attacks glass. Consequently, fluorides of alkali metals produce basic solutions. For example, a 1 M solution of NaF in water has a pH of 8.59 compared to a 1 M solution of NaOH, a strong base, which has a pH of 14.00.[2]
2
Fluorine
3
Isotopes Although fluorine (F) has multiple isotopes, only one of these isotopes (F-19) is stable, and the others have short half-lives and are not found in nature. Fluorine is thus a mononuclidic element. The nuclide 18F is the radionuclide of fluorine with the longest half life (about 110 minutes = almost 2 hours), and commercially is an important source of positrons-- finding its major use in positron emission tomography scanning.
Applications Elemental fluorine, F2, is mainly used for the production of two compounds of commercial interest, uranium hexafluoride and sulfur hexafluoride.[3] Industrial use of fluorine-containing compounds: • Atomic fluorine and molecular fluorine are used for plasma etching in semiconductor manufacturing, flat panel display production and MEMS (microelectromechanical systems) fabrication.[4] Xenon difluoride is also used for this last purpose. • Hydrofluoric acid (chemical formula HF) is used to etch glass in light bulbs and other products. • Tetrafluoroethylene and perfluorooctanoic acid (PFOA) are directly used in the production of low friction plastics such as Teflon (or polytetrafluoroethylene). • Fluorine is used indirectly in the production of halons such as freon. • Along with some of its compounds, fluorine is used in the production of pure uranium from uranium hexafluoride and in the synthesis of numerous commercial fluorochemicals, including vitally important pharmaceuticals, agrochemical compounds, lubricants, and textiles. • Fluorochlorohydrocarbons are used extensively in air conditioning and in refrigeration. Chlorofluorocarbons have been banned for these applications because they contribute to ozone destruction and the ozone hole. Interestingly, since it is chlorine and bromine radicals which harm the ozone layer, not fluorine, compounds which do not contain chlorine or bromine but contain only fluorine, carbon and hydrogen (called hydrofluorocarbons) are not on the United States Environmental Protection Agency list of ozone-depleting substances,[5] and have been widely used as replacements for the chlorine- and bromine-containing fluorocarbons. Hydrofluorocarbons do have a greenhouse effect, but a small one compared with carbon dioxide and methane. • Sodium hexafluoroaluminate (cryolite), is used in the electrolysis of aluminium. • In much higher concentrations, sodium fluoride has been used as an insecticide, especially against cockroaches. • Fluorides have been used in the past to help molten metal flow, hence the name. • Some researchers including US space scientists in the early 1960s have studied elemental fluorine gas as a possible rocket propellant due to its exceptionally high specific impulse. The experiments failed because fluorine proved difficult to handle, and its combustion products proved extremely toxic and corrosive. • Compounds of fluorine such as fluoropolymers, potassium fluoride and cryolite are utilized in applications such as anti-reflective coatings and dichroic mirrors on account of their unusually low refractive index.
Fluorine
Dental and medical uses • Inorganic compounds of fluoride, including sodium fluoride (NaF), stannous fluoride (SnF2) and sodium MFP, are used in toothpaste to prevent dental cavities. These or related compounds are also added to some municipal water supplies, a process called water fluoridation, although the practice has remained controversial since its beginnings in 1945. • Many important agents for general anesthesia such as sevoflurane, desflurane, and isoflurane are hydrofluorocarbon derivatives. • The fluorinated antiinflammatories dexamethasone and triamcinolone are among the most potent of the synthetic corticosteroids class of drugs.[6] • Fludrocortisone ("Florinef") is one of the most common mineralocorticoids, a class of drugs which mimics the actions of aldosterone. • Fluconazole is a triazole antifungal drug used in the treatment and prevention of superficial and systemic fungal infections. • Fluoroquinolones are a family of broad-spectrum antibiotics. • SSRI antidepressants, except in a few instances, are fluorinated molecules. These include citalopram, escitalopram oxalate, fluoxetine, fluvoxamine maleate, and paroxetine. A notable exception is sertraline. Because of the difficulty of biological systems in dealing with metabolism of fluorinated molecules, fluorinated antibiotics and antidepressants are among the major fluorinated organics found in treated city sewage and wastewater. • Compounds containing 18F, a radioactive isotope that emits positrons, are often used in positron emission tomography, because its half-life of 110 minutes is long by the standards of positron-emitters. One such species is fluorodeoxyglucose.
Chemistry of fluorine Fluorine forms a variety of very different compounds, owing to its small atomic size and covalent behavior. Elemental fluorine is a dangerously powerful oxidant, reflecting the extreme electronegativity of fluorine. Hydrofluoric acid is extremely dangerous, whereas in synthetic drugs incorporating an aromatic ring (e.g. flumazenil), fluorine is used to help prevent toxication or to delay metabolism. The fluoride ion is basic, therefore hydrofluoric acid is a weak acid in water solution. However, water is not an inert solvent in this case: when less basic solvents such as anhydrous acetic acid are used, hydrofluoric acid is the strongest of the hydrohalogenic acids. Also, owing to the basicity of the fluoride ion, soluble fluorides give basic water solutions. The fluoride ion is a Lewis base, and has a high affinity to certain elements such as calcium and silicon. For example, deprotection of silicon protecting groups is achieved with a fluoride. The fluoride ion is poisonous. Fluorine as a freely reacting oxidant gives the strongest oxidants known. The reactivity of fluorine toward the noble gas xenon was first reported by Neil Bartlett in 1962. Fluorides of krypton and radon have also been prepared. Argon fluorohydride has been observed at cryogenic temperatures. The carbon-fluoride bond is covalent and very stable. The use of a fluorocarbon polymer, poly(tetrafluoroethene) or Teflon, is an example: it is thermostable and waterproof enough to be used in frying pans. Organofluorines may be safely used in applications such as drugs, without the risk of release of toxic fluoride. In synthetic drugs, toxication can be prevented.
4
Fluorine
5
For example, an aromatic ring is useful but presents a safety problem: enzymes in the body metabolize some of them into poisonous epoxides. When the para position is substituted with fluorine, the aromatic ring is protected and epoxide is no longer produced. The substitution of fluorine for hydrogen in organic compounds offers a very large number of compounds. An estimated fifth of pharmaceutical compounds and 30% of agrochemical compounds contain fluorine.[7] The -CF3 and -OCF3 moieties provide further variation, and more recently the -SF5 group.[8]
Fluorite (CaF2) crystals
Production Industrial production of fluorine entails the electrolysis of hydrogen fluoride in the presence of potassium fluoride. This method is based on the pioneering studies by Moissan (see below). Fluorine gas forms at the anode, and hydrogen gas at the cathode. Under these conditions, the potassium fluoride (KF) converts to potassium bifluoride (KHF2), which is the actual electrolyte, This potassium bifluoride aids electrolysis by greatly increasing the electrical conductivity of the solution.
Fluorine cell room at F2 Chemicals Ltd, Preston, UK
HF + KF → KHF2 2 KHF2 → 2 KF + H2 + F2 The HF required for the electrolysis is obtained as a byproduct of the production of phosphoric acid. Phosphate-containing minerals contain significant amounts of calcium fluorides, such as fluorite. Upon treatment with sulfuric acid, these minerals release hydrogen fluoride:
Fluorine
6 CaF2 + H2SO4 → 2 HF + CaSO4
In 1986, when preparing for a conference to celebrate the 100th anniversary of the discovery of fluorine, Karl Christe discovered a purely chemical preparation involving the reaction of solutions in anhydrous HF, K2MnF6, and SbF5 at 150 °C:[9] K2MnF6 + 2 SbF5 → 2 KSbF6 + MnF3 + ½ F2 Though not a practical synthesis on the large scale, this report demonstrates that electrolysis is not the sole route to the element.
History The mineral fluorspar (also called fluorite), consisting mainly of calcium fluoride, was described in 1530 by Georgius Agricola for its use as a flux.[10] Fluxes are used to promote the fusion of metals or minerals. The etymology of the element's name reflects its history: Fluorine pronounced /ˈflʊəriːn/, /ˈflʊərɨn/, or commonly /ˈflɔr-/; from Latin: fluere, meaning "to flow". In 1670 Schwanhard found that glass was etched when it was exposed to fluorspar that had been treated with acid. Carl Wilhelm Scheele and many later researchers, including Humphry Davy, Caroline Menard, Gay-Lussac, Antoine Lavoisier, and Louis Thenard all would experiment with hydrofluoric acid, easily obtained by treating fluorite with concentrated sulfuric acid. Owing to its extreme reactivity, elemental fluorine was not isolated until many years after the characterization of fluorite. Progress in isolating elemental fluorine was slowed because it could only be prepared electrolytically and even then under stringent conditions since the gas attacks many materials. In 1886, the isolation of elemental fluorine was reported by Henri Moissan after almost 74 years of effort by other chemists.[11] The generation of elemental fluorine from hydrofluoric acid is exceptionally dangerous, killing or blinding several scientists who attempted early experiments on this halogen. These individuals came to be referred to as "fluorine martyrs".[12] For Moissan, it earned him the 1906 Nobel Prize in chemistry.[13] The first large-scale production of fluorine was undertaken in support of the Manhattan project, where the compound uranium hexafluoride (UF6) had been selected as the form of uranium that would allow separation of its 235U and 238U isotopes. Today both the gaseous diffusion process and the gas centrifuge process use gaseous UF6 to produce enriched uranium for nuclear power applications. In the Manhattan Project, it was found that UF6 decomposed into UF4 and F2. The corrosion problem due to the F2 was eventually solved by electrolytically coating all UF6 carrying piping with nickel metal, which forms a nickel difluoride that is not attacked by fluorine. Joints and flexible parts were made from teflon, then a very recently discovered fluorocarbon plastic which is also not attacked by F2.
Biological role Though F2 is too reactive to have any natural biological role, fluorine is incorporated into compounds with biological activity. Naturally occurring organofluorine compounds are rare, the most notable example is fluoroacetate, which functions as a plant defence against herbivores in at least 40 plants in Australia, Brazil and Africa.[14] The enzyme adenosyl-fluoride synthase catalyzes the formation of 5'-deoxy-5'-fluoroadenosine. Fluorine is not an essential nutrient, but its importance in preventing tooth decay is well-recognized.[15] The effect is predominantly topical, although prior to 1981 it was
Fluorine
7
considered primarily systemic (occurring through ingestion).[16]
Precautions Elemental fluorine Elemental fluorine (fluorine gas) is a highly toxic, corrosive oxidant, which can cause ignition of organic material. Fluorine gas has a characteristic pungent odor that is detectable in concentrations as low as 20 ppb. As it is so reactive, all materials of construction must be carefully selected and metal surfaces must be passivated.
Fluoride ion Fluoride ions are toxic: the lethal dose of sodium fluoride for a 70 kg human is estimated at 5–10 g.[17]
Hydrogen fluoride and hydrofluoric acid Hydrogen fluoride and hydrofluoric acid are dangerous, far more so than the related hydrochloric acid, because undissociated molecular HF penetrates the skin and biological membranes, causing deep and painless burns. The free fluoride, once released from HF in dissociation, also is capable of chelating calcium ion to the point of causing death by cardiac dysrhythmia. Burns with areas larger than 25 square inches (160 cm2) have the potential to cause serious systemic toxicity.[18]
Organofluorines Organofluorines are naturally rare compounds. They can be nontoxic (perflubron and perfluorodecalin) or highly toxic (perfluoroisobutylene and fluoroacetic acid). Many pharmacuticals are organofluorines, such as the anti-cancer fluorouracil. Perfluorooctanesulfonic acid (PFOS) is a persistent organic pollutant.
See also • • • •
Fluorocarbon Isotopes of fluorine Halide minerals Water fluoridation
References • Los Alamos National Laboratory – Fluorine
[19]
External links • • • •
WebElements.com – Fluorine [20] It's Elemental – Fluorine [21] Picture of liquid fluorine – chemie-master.de Chemsoc.org [23]
[22]
Fluorine
References [1] Theodore Gray. " Real visible fluorine (http:/ / theodoregray. com/ PeriodicTable/ Samples/ 009. 5/ index. s12. html)". The Wooden Periodic Table. . [2] " pKa's of Inorganic and Oxo-Acids (http:/ / www2. lsdiv. harvard. edu/ labs/ evans/ pdf/ evans_pKa_table. pdf)". Evans Group. . Retrieved 2008-11-29. [3] M. Jaccaud, R. Faron, D. Devilliers, R. Romano (2005). Fluorine, in Ullmann’s Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. ISBN 3527310975. [4] Leonel R Arana, Nuria de Mas, Raymond Schmidt, Aleksander J Franz, Martin A Schmidt and Klavs F Jensen (2007). "Isotropic etching of silicon in fluorine gas for MEMS micromachining". J. Micromech. Microeng. 17: 384. doi: 10.1088/0960-1317/17/2/026 (http:/ / dx. doi. org/ 10. 1088/ 0960-1317/ 17/ 2/ 026). [5] " Class I Ozone-Depleting Substances (http:/ / www. epa. gov/ ozone/ ods. html)". Ozone Depletion. U.S. Environmental Protection Agency. . [6] Steve S Lim. " eMedicine - Corticosteroid-Induced Myopathy (http:/ / www. emedicine. com/ pmr/ topic35. htm)". . [7] " Fluorine's treasure trove (http:/ / www. icis. com/ Articles/ 2006/ 09/ 30/ 2016413/ fluorines-treasure-trove. html)". ICIS news. 2006-10-02. . Retrieved 2008-11-29. [8] Bernhard Stump, Christian Eberle, W. Bernd Schweizer, Marcel Kaiser, Reto Brun, R. Luise Krauth-Siegel, Dieter Lentz, François Diederich (2009). "Pentafluorosulfanyl as a Novel Building Block for Enzyme Inhibitors: Trypanothione Reductase Inhibition and Antiprotozoal Activities of Diarylamines". ChemBioChem 10: 79. doi: 10.1002/cbic.200800565 (http:/ / dx. doi. org/ 10. 1002/ cbic. 200800565). [9] K. Christe (1986). "Chemical synthesis of elemental fluorine". Inorg. Chem. 25: 3721–3724. doi: 10.1021/ic00241a001 (http:/ / dx. doi. org/ 10. 1021/ ic00241a001). [10] " Discovery of fluorine (http:/ / www. fluoride-history. de/ fluorine. htm)". Fluoride History. . [11] H. Moissan (1886). " Action d'un courant électrique sur l'acide fluorhydrique anhydre (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k3058f/ f1541. chemindefer)". Comptes rendus hebdomadaires des séances de l'Académie des sciences 102: 1543–1544. . [12] Richard D. Duncan. (2008). Elements of faith : faith facts and learning lessons from the periodic table (http:/ / books. google. de/ books?id=kgVAlzGXx6oC). Green Forest, Ark.: Master Books. p. 22. ISBN 9780890515471. . [13] " The Nobel Prize in Chemistry 1906 (http:/ / nobelprize. org/ nobel_prizes/ chemistry/ laureates/ 1906/ )". Nobelprize.org. . Retrieved 2009-07-07. [14] Proudfoot AT, Bradberry SM, Vale JA (2006). "Sodium fluoroacetate poisoning". Toxicol Rev 25 (4): 213–9. doi: 10.2165/00139709-200625040-00002 (http:/ / dx. doi. org/ 10. 2165/ 00139709-200625040-00002). PMID 17288493. [15] Olivares M and Uauy R (2004). " Essential nutrients in drinking-water (Draft) (http:/ / www. who. int/ water_sanitation_health/ dwq/ en/ nutoverview. pdf)". WHO. . Retrieved 2008-12-30. [16] Pizzo G, Piscopo MR, Pizzo I, Giuliana G (September 2007). "Community water fluoridation and caries prevention: a critical review". Clin Oral Investig 11 (3): 189–93. doi: 10.1007/s00784-007-0111-6 (http:/ / dx. doi. org/ 10. 1007/ s00784-007-0111-6). PMID 17333303. [17] Aigueperse, Jean; Paul Mollard, Didier Devilliers, Marius Chemla, Robert Faron, Renée Romano, Jean Pierre Cuer (2005), "Fluorine Compounds, Inorganic", in Ullmann, Encyclopedia of Industrial Chemistry, Weinheim: Wiley-VCH [18] " Recommended Medical Treatment for Hydrofluoric Acid Exposure (http:/ / www51. honeywell. com/ sm/ hfacid/ common/ documents/ HF_medical_book. pdf)" (PDF). Honeywell Specialty Materials. . Retrieved 2009-05-06. [19] [20] [21] [22] [23]
http:/ / periodic. lanl. gov/ elements/ 9. html http:/ / www. webelements. com/ fluorine/ http:/ / education. jlab. org/ itselemental/ ele009. html http:/ / www. chemie-master. de/ pse/ pse. php?modul=F http:/ / www. chemsoc. org/ viselements/ pages/ fluorine. html
8
Article Sources and Contributors
Article Sources and Contributors Fluorine Source: http://en.wikipedia.org/w/index.php?oldid=306561968 Contributors: 12dstring, A. Carty, A3r0, AAAAA, ACSE, Ace of Spades IV, Aeluwas, Aerion, Ahoerstemeier, Alansohn, AllHailZeppelin, Amazon10x, Andres, Andrewa, Antandrus, Anthony Appleyard, Antidisestablishmentarinism, Antonio Lopez, Archimerged, AtheWeatherman, AxelBoldt, Axeman89, Az7997, Beantwo, Beetstra, Benbest, Benjah-bmm27, BlackIvy88, BlastOButter42, Blazersmel, BlueEarth, Bobo192, Bomac, Brianga, Brianski, Brighterorange, Bryan Derksen, Bubba Joe, Buckslayer, ByAppointmentTo, CWii, CYD, Camembert, Can't sleep, clown will eat me, Capricorn42, Carboxen, CardinalDan, Carey Evans, Carnildo, Causesobad, Centrx, Ched Davis, ChemNerd, Chris Dybala, Chris G, ChrisHamburg, Christian List, Circeus, ClanCC, Clutch, Cmichael, Colbuckshot, Conversion script, Cosmium, Cxz111, DMacks, Dajwilkinson, Dalesgay, DanielCD, Darrenr99, Darrien, David Latapie, Davidhorman, Deli nk, Delirium, Delldot, DennyColt, DerHexer, Discospinster, Doulos Christos, DrBob, Drini, Dukeofomnium, Dwmyers, Dycedarg, Dylan Lake, EGGS, EL Willy, Ed Poor, EdBever, Edgar181, Edgepedia, Edsanville, Ee79, Ekrub-ntyh, El C, Elassint, Eldin raigmore, Elerium, Eman120494, Emperorbma, Enviroboy, Epbr123, Eric-Wester, Eric119, Erik Zachte, Esrever, F2Andy, Femto, Fivemack, Flamingrok3, Flink the blind hemophiliac, Flyguy649, FreplySpang, Frosty0814snowman, Gene Nygaard, Giftlite, Gman124, Grendelkhan, Groucho NL, Ground Zero, Gurch, Guyzero, Hadal, Hak-kâ-ngìn, Halokid211, HappyCamper, Hawaiian717, Hdt83, Heron, HexaChord, Hike395, Hurricane Devon, II MusLiM HyBRiD II, IRP, Icairns, Icewedge, ImperfectlyInformed, Iridescent, Itub, Ixfd64, J.delanoy, Jack, Jacobbirdy128, James086, Jaraalbe, Jclemens, Jcook56050, Jessemerriman, Jjasi, JoanneB, John, Johnbrownsbody, Jose77, Jrockley, Jushi, Karl-Henner, Karlhahn, Kbdank71, Keenan Pepper, Kelly Martin, Kenken71, Kilo-Lima, Kkmurray, KnowledgeOfSelf, Koyaanis Qatsi, Kpjas, Krm500, Ktsquare, Kukini, Kurykh, Kwamikagami, LOL, Laguna72, Lankiveil, Leonard G., Leuko, Leytonwd, Lmbstl, LorenzoB, LouScheffer, Luna Santin, MZMcBride, Malbi, Mandor, MarcoTolo, Mark Ryan, MarkSweep, Marnanel, MartinHarper, Masterx, Materialscientist, Matnkat, Matthewmayer, Mav, Maximillion Pegasus, Maxis ftw, MchlWngr, Meeples, Megaboz, Mentisock, Mgimpel, Mikez, Minesweeper, Mo0, MoogleDan, Mr Stephen, Mr.Z-man, Mrholybrain, Musser, Mwoolf, Mww113, Mxn, N4nojohn, NPChristmas, Nakon, Nanogene, NawlinWiki, Neil916, Nergaal, Neurolysis, Nick, Nihiltres, Nonagonal Spider, NuclearWarfare, Obiwanskywalker, Oliphaunt, Oxymoron83, Patstuart, Paul August, Pharaoh of the Wizards, Philip Trueman, Piano non troppo, PierreAbbat, PigFlu Oink, Pishogue, Plumbago, PoliteCarbide, Polonium, Poolkris, Poor Yorick, PopUpPirate, PrestonH, Puchiko, Quintote, RTC, Ratherhaveaheart, Raul654, Razorflame, Remember, Reusche, Revived, RexNL, Reza kalani, Richnotts, Rifleman 82, Rmhermen, RobertAustin, Roberta F., Robertb-dc, Romanm, Rrburke, Rursus, SJP, Sam Korn, Saperaud, Sbharris, Sceptre, Schneelocke, Sengkang, Sharkface217, Shawn in Montreal, Sheitan, Shimmin, Shirifan, Shootbamboo, Sietse Snel, SigmaEpsilon, Sk8er5000, Skatebiker, Sl, Smokefoot, Snezzy, Snowolf, Soliloquial, Solipsist, Someone else, Sorfane, SorrryCharlie, Spellmaster, Spiffytease45, Squids and Chips, Starom, StaticGull, Stephen Gilbert, Stephenb, Stifynsemons, Stone, Suisui, Sullivan.t.j, Sunborn, Synchronism, T-borg, TeaDrinker, Tengfred, Tetracube, That Guy, From That Show!, Thaurisil, The Rambling Man, The sunder king, TheBendster, Thehelpfulone, Thingg, Thricecube, Tide rolls, Tim Starling, Tloser, Tohd8BohaithuGh1, Tom harrison, Travis.Thurston, Tsunaminoai, Twang, Velvetron, Vera Cruz, Vhyntsze, Viktor-viking, Vlad4599, Vsmith, Vssun, Vuo, Warut, Washburnmav, Watch37264, Wes!, Wikiscrewcumdumpsta, Wknight94, Wtmitchell, Wyllium, Xaosflux, Xenophon777, Yamamoto Ichiro, Yyy, Zach4636, Ziyaddd, Τις, 840 anonymous edits
Image Sources, Licenses and Contributors file:cubic.svg Source: http://en.wikipedia.org/w/index.php?title=File:Cubic.svg License: GNU Free Documentation License Contributors: User:Stannered file:Electron shell 009 Fluorine.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_009_Fluorine.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 Image:Fluorite crystals 270x444.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Fluorite_crystals_270x444.jpg License: GNU Free Documentation License Contributors: HereToHelp, Jurema Oliveira, MushiHoshiIshi, Ra'ike, Saperaud Image:Fluorine cell room.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Fluorine_cell_room.jpg License: Public Domain Contributors: F2 Chemicals Ltd (AK Joel)
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9
Neon
1
Neon fluorine ← neon → sodiumHe ↑ Ne ↓ Ar
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HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
10Ne Periodic table
Appearance colorless gas General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensityMelting pointBoiling pointTriple pointCritical pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
12
13
15
18
21
27
Atomic properties Oxidation states Ionization energies (more) 2nd: 3952.3 kJ·mol−1 3rd: 6122 kJ·mol−1Covalent radiusVan der Waals radius Miscellaneous Crystal structureMagnetic orderingThermal conductivitySpeed of soundBulk modulusCAS registry number Most stable isotopes Main article: Isotopes of neon iso
N.A.
half-life
20
90.48%
20
21
0.27%
21
22
9.25%
22
Ne Ne Ne
DM
Ne is stable with 10 neutron Ne is stable with 11 neutron Ne is stable with 12 neutron
DE (MeV)
DP
Neon neon, Ne, 10 noble gases 18, 2, p20.1797(6) g·mol−1 1s2 2s2 2p6 2, 8 (Image) gas (0 °C, 101.325 kPa) 0.9002 g/L 24.56 K,-248.59 °C,-415.46 °F 27.07 K,-246.08 °C,-410.94 °F 24.5561 K (-249°C), 43[1] [2] kPa 44.4 K, 2.76 MPa 0.335 kJ·mol−1 1.71 kJ·mol−1 (25 °C) −1 −1 −1 20.786 J·mol ·K no data 1st: 2080.7 kJ·mol 58 pm 154 pm face-centered cubic diamagnetic[3] (300 K) 49.1x10-3 W·m−1·K−1 (gas, 0 °C) 435 m/s 654 GPa 7440-01-9 Neon (pronounced /ˈniːɒn/) is the chemical element that has the symbol Ne and atomic number 10. Although a very common element in the universe, it is rare on Earth. A colorless, inert noble gas under standard conditions, neon gives a distinct reddish-orange glow when used in discharge tubes and neon lamps.[4] [5] It is commercially extracted from air, in which it is found in trace amounts.
History Neon (Greek νέον (neon) meaning "new one") was discovered in 1898 by Scottish chemist Sir William Ramsay (1852–1916) English chemist Morris W. Travers (1872–1961) in London, England.[6] Neon was discovered when Ramsay chilled a sample of the atmosphere until it became a liquid, then warmed the liquid and captured the gases as they boiled off. The three gases that boiled off were krypton, xenon, and neon.[7] In December 1910, French engineer Georges Claude made a lamp from an electrified tube of neon gas. On January 19, 1915, Claude began selling his tubes to U.S. companies; the Packard car dealership in Los Angeles was one of the first to buy it.[8]
Isotopes Neon has three stable isotopes: 20Ne (90.48%), 21Ne (0.27%) and 22Ne (9.25%). 21Ne and 22 Ne are nucleogenic and their variations are well understood. In contrast, 20Ne is not known to be nucleogenic and the causes of its variation in the Earth have been hotly debated. The principal nuclear reactions which generate neon isotopes are neutron emission, alpha decay reactions on 24Mg and 25Mg, which produce 21Ne and 22Ne, respectively. The alpha particles are derived from uranium-series decay chains, while the neutrons are mostly produced by secondary reactions from alpha particles. The net result yields a trend towards lower 20Ne/22Ne and higher 21Ne/22Ne ratios observed in uranium-rich rocks such as granites. Isotopic analysis of exposed terrestrial rocks has demonstrated the cosmogenic production of 21Ne. This isotope is generated by spallation reactions on magnesium, sodium, silicon, and aluminium. By analyzing all three isotopes, the cosmogenic component can be resolved from magmatic neon and nucleogenic neon. This suggests that neon will be a useful tool in determining cosmic exposure ages of surficial rocks and meteorites.[9] Similar to xenon, neon content observed in samples of volcanic gases are enriched in 20Ne, as well as nucleogenic 21Ne, relative to 22Ne content. The neon isotopic content of these mantle-derived samples represent a non-atmospheric source of neon. The 20Ne-enriched components are attributed to exotic primordial rare gas components in the Earth, possibly representing solar neon. Elevated 20Ne abundances are found in diamonds, further suggesting a solar neon reservoir in the Earth.[10]
2
Neon
Characteristics Neon is the second-lightest noble gas. It glows reddish-orange in a vacuum discharge tube. According to recent studies, neon is the least reactive noble gas and thus the least reactive of all elements[11] . Also, neon has the narrowest liquid range of any element: from 24.55 K to 27.05 K (-248.45 °C to −245.95 °C, or −415.21 °F to −410.71 °F). It has over 40 times the refrigerating capacity of liquid helium and three times that of liquid hydrogen (on a per unit volume basis).[12] In most applications it is a less expensive refrigerant than helium.[13] Neon plasma has the most intense light discharge at normal voltages and currents of all the noble gases. The average colour of this light to the human eye is red-orange Spectrum of neon with ultraviolet lines (at left) and infrared (at right) shown in white due to many lines in this range; it also contains a strong green line which is hidden, unless the visual components are dispersed by a spectroscope.[14] Two quite different kinds of neon lights are in common use. Glow-discharge lamps are typically tiny, and often designed to operate at 120 volts; they are widely used as power-on indicators and in circuit-testing equipment. Neon signs and other arc-discharge devices operate instead at high voltages, often 3–15 kilovolts (3,000–15,000 volts); they can be made into (often bent) tubes a few meters long.
Occurrence Neon is actually abundant on a universal scale: the fifth most abundant chemical element in the universe by mass, after hydrogen, helium, oxygen, and carbon (see chemical element). Its relative rarity on Earth, like that of helium, is due to its relative lightness and chemical inertness, both properties keeping it from being trapped in the condensing gas and dust clouds of the formation of smaller and warmer solid planets like Earth. Neon is monatomic, making it lighter than the molecules of diatomic nitrogen and oxygen which form the bulk of Earth's atmosphere; a balloon filled with neon will rise up into the air, albeit more slowly than a helium balloon.[15] Mass abundance in the universe is about 1 part in 750 and in the Sun and presumably in the proto-solar system nebula, about 1 part in 600. The Galileo spacecraft atmospheric entry probe found that even in the upper atmosphere of Jupiter, neon is reduced by about a factor of 10, to 1 part in 6,000 by mass. This may indicate that even the ice-planetesmals which brought neon into Jupiter from the outer solar system, formed in a region which was too warm for them to have kept their neon (abundances of heavier inert gases on Jupiter are several times that found in the Sun).[16] Neon is a monatomic gas at standard conditions. Neon is rare on Earth, found in the Earth's atmosphere at 1 part in 65,000 (by volume) or 1 part in 83,000 by mass. It is industrially produced by cryogenic fractional distillation of liquefied air.[12]
3
Neon
4
Applications Neon is often used in signs and produces an unmistakable bright reddish-orange light. Although still referred to as "neon", all other colours are generated with the other Noble Gases or by many colours of fluorescent lighting. Neon is used in vacuum tubes, high-voltage indicators, lightning arrestors, wave meter tubes, television tubes, and helium-neon lasers. Liquefied neon is commercially used as a cryogenic refrigerant in applications not A neon sign in the shape of its name. requiring the lower temperature range attainable with more extreme liquid helium refrigeration. Liquid neon is actually quite expensive, and nearly impossible to obtain in small quantities for laboratory tests. For small quantities, liquid neon can be >55x more expensive than liquid helium. The driver for expense is actually rarity of the gas, not the liquefaction process. The triple point temperature of Neon (24.5561 K) is a defining fixed point in the International Temperature Scale of 1990.[1]
Compounds Neon is the first p-block noble gas. Theoretically neon is the least reactive of all noble gases (including helium which produces a metastable compound HHeF), and therefore generally considered to be inert. The calculated bond energies of neon with noble metals, hydrogen, beryllium and boron are lesser than that of helium or any other noble gas. No true compounds including the neutral compounds of neon are known. However, the ions Ne+, (NeAr)+, (NeH)+, and (HeNe+) have been observed from optical and mass spectrometric studies, and there are some unverified reports of an unstable hydrate.[12]
See also • Expansion ratio • Neon sign • Neon lamp
External links • • • • • •
WebElements.com – Neon [17] It's Elemental – Neon [18] Computational Chemistry Wiki [19] USGS Periodic Table - Neon [20] Atomic Spectrum of Neon [21] Neon Museum, Las Vegas [22]
Neon
5
References [1] Preston-Thomas, H. (1990). " The International Temperature Scale of 1990 (ITS-90) (http:/ / www. bipm. org/ en/ publications/ its-90. html)". Metrologia 27: 3-10. . [2] "Section 4, Properties of the Elements and Inorganic Compounds; Melting, boiling, triple, and critical temperatures of the elements". CRC Handbook of Chemistry and Physics (85th edition ed.). Boca Raton, Florida: CRC Press. 2005. [3] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [4] Harold P Coyle (2001). Project STAR: The Universe in Your Hands (http:/ / books. google. com/ books?id=KwTzo4GMlewC& pg=PA127). Kendall Hunt. ISBN 9780787267636. . [5] Kohtaro Kohmoto (1999). " Phosphors for lamps (http:/ / books. google. com/ books?id=lWlcJEDukRIC& pg=PA380)". in Shigeo Shionoya and William M. Yen. Phosphor Handbook. CRC Press. ISBN 9780849375606. . [6] William Ramsay, Morris W. Travers (1898). " On the Companions of Argon (http:/ / www. jstor. org/ pss/ 116011)". Proceedings of the Royal Society of London 63: 437–440. doi: 10.1098/rspl.1898.0057 (http:/ / dx. doi. org/ 10. 1098/ rspl. 1898. 0057). . [7] " Neon: History (http:/ / nautilus. fis. uc. pt/ st2. 5/ scenes-e/ elem/ e01000. html)". Softciências. . Retrieved 2007-02-27. [8] " Neon: A Brief History (http:/ / nymag. com/ shopping/ features/ 41814/ )". New York Magazine. . Retrieved 2008-05-20. [9] " Neon: Isotopes (http:/ / nautilus. fis. uc. pt/ st2. 5/ scenes-e/ elem/ e01093. html)". Softciências. . Retrieved 2007-02-27. [10] Anderson, Don L.. " Helium, Neon & Argon (http:/ / www. mantleplumes. org/ Ne. html)". Mantleplumes.org. . Retrieved 2006-07-02. [11] Errol G. Lewars (2008). " Modelling Marvels (http:/ / books. google. co. in/ books?id=whdw2qlXjD0C& pg)". Springer. . [12] C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81th edition (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elements. pdf). CRC press. p. 19. ISBN 0849304814. . [13] " NASSMC: News Bulletin (http:/ / www. nassmc. org/ bulletin/ dec05bulletin. html#table)". 30. . Retrieved 2007-03-05. [14] " Plasma (http:/ / www. electricalfun. com/ plasma. htm)". . Retrieved 2007-03-05. [15] R. Gallagher and P. Ingram (2001). Chemistry for Higher Tier (http:/ / books. google. com/ books?id=SJtWSy69eVsC& pg=PA96). University Press. ISBN 9780199148172. . [16] Morse, David (26). " Galileo Probe Science Result (http:/ / www2. jpl. nasa. gov/ sl9/ gll38. html)". Galileo Project. . Retrieved 2007-02-27. [17] [18] [19] [20] [21] [22]
http:/ / www. webelements. com/ neon/ http:/ / education. jlab. org/ itselemental/ ele010. html http:/ / www. compchemwiki. org/ index. php?title=Neon http:/ / wwwrcamnl. wr. usgs. gov/ isoig/ period/ ne_iig. html http:/ / hyperphysics. phy-astr. gsu. edu/ Hbase/ quantum/ atspect2. html http:/ / www. neonmuseum. org/
Article Sources and Contributors
Article Sources and Contributors Neon Source: http://en.wikipedia.org/w/index.php?oldid=307518699 Contributors: 000lynx, 05clareb, A new name 2008, A2Kafir, ABF, AEMoreira042281, Aadgray, Abrech, Adashiel, Adrian.benko, Aeluwas, Aeon1006, Ahoerstemeier, Aitias, Alansohn, Ale jrb, Aleenf1, Alexanderveringa, Alphachimp, Amicon, Anclation, Andre Engels, AndreasJS, Andres, Anonymous Dissident, Anoop.m, Antandrus, Anthony, Aperram, Arakunem, Archanamiya, Archimerged, ArglebargleIV, Arjun01, AtheWeatherman, BRG, Badocter, Baronnet, Basawala, Batmanbb, BeaverMonkey, Beetstra, Berkunt, Bexxx x3, Big Bird, Black-Velvet, BlastOButter42, BlueEarth, Bobbyjoe58963, Bobo192, Bogey97, Bomac, Briscanator, Bubba hotep, BunsenH, Buttchug, Bye bye now, CYD, Calair, Camw, Can't sleep, clown will eat me, CanisRufus, Capricorn42, Capybara21, CardinalDan, Carnildo, Casper2k3, Casull, Catgut, ChemNerd, Chitrapa, Chris Dybala, Chrislk02, Chuggo, Chun-hian, Closedmouth, Cloud Strife, Conversion script, Cool Blue, Cosmium, Cratbro, Croat Canuck, Cryptic, Cryptic C62, Crystal whacker, Ctjf83, Cureden, DMacks, DSRH, Darkakatsuki, Darrien, Darth Panda, Dataphile, Dave6, David Latapie, David.Gross, Davidkazuhiro, Dbtfz, DeadEyeArrow, Deglr6328, Deli nk, Delldot, Dendlai, Deon, DerHexer, Derek.cashman, Dgrant, Dicklyon, Diligent Terrier, Dina, Discospinster, Doctorblaze, Doctorevil64, Doctormatt, Doulos Christos, DrBob, Dreadstar, Drini, Dupz, EL Willy, ESkog, Edgar181, Edsanville, Electricmic, Electron9, Elipongo, Emperorbma, Eng02019, Enviroboy, Epbr123, Epo, Equendil, Eric119, Erik Zachte, Euicho, Evercat, Everyking, Farosdaughter, Fashiondebz, FastLizard4, Feline1, Femto, Fffhfhfh, Fieldday-sunday, Finlay McWalter, Fonzy, FrankCostanza, Frankie0607, Fredpfi, FreplySpang, Frozenevolution, Gamer007, Garden, Gary King, George The Dragon, Giftlite, Gilliam, Gimmetrow, Gogo Dodo, Goodnightmush, Gracenotes, GraemeL, Grendelkhan, Gwernol, Gzkn, Hak-kâ-ngìn, Hall Monitor, Helge Skjeveland, Hellbus, HenryLi, Herbee, Heron, Hersfold, Hyphz, Iago4096, Icairns, Igksb, Ilovetractors, Independent Journalist, Indon, Indosauros, Instinct, Iridescent, Ixfd64, J Di, J.delanoy, JForget, Jack312, Jake Wartenberg, Jaknouse, James086, Jan.Smolik, Jaraalbe, Jayeeararwhy, Jcw69, Jdrewitt, Jdurg, Jeanettedugas, Jemijohn, Jj137, Joanjoc, Jobe6, Johann Wolfgang, John, JohnWittle, Jojhutton, Jondel, Jose77, Jumbuck, K10102898, KP Shadowww, Karl-Henner, Karlhahn, Keilana, Kelly Martin, Ker-Jar Song, Kerowyn, Kookid654, Kostisl, Kowey, Koyaanis Qatsi, Kpjas, Krashlandon, Kukini, Kuru, Kwamikagami, LA2, LarryMorseDCOhio, LeaveSleaves, Letstalk, Lexor, LibLord, Lightdarkness, Lightmouse, Ll11812842, Looxix, Lop7685, Loren.wilton, Lost tourist, LostArtilleryman, Luckas Blade, Luke Green, Lupin, MZMcBride, Mada48, Maerk, Majorclanger, Malcolm, ManoaChild, Marnanel, Massimo Catarinella, Master Jay, Materialscientist, Matt Yeager, Mav, McSly, Megaboz, Meno25, Mentifisto, Mgimpel, Michaelas10, Michbich, 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6
Sodium
1
Sodium neon ← sodium → magnesiumLi ↑ Na ↓ K
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
11Na Periodic table
Appearance silvery white metallic
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointCritical pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
554
617
697
802
946
1153
Sodium
2
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 4562 kJ·mol−1 3rd: 6910.3 kJ·mol−1Atomic radiusCovalent radiusVan der Waals radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusMohs hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of sodium iso 22
Na
N.A. syn
half-life 2.602 y
DM β+→γ
ε→γ
β+ 23
Na
100%
DE (MeV)
DP
0.5454
22
[1] 1.27453(2)
22
-
22
1.27453(2)
22
1.8200
22
Ne* Ne Ne* Ne Ne
23
Na is stable with 12 neutron
sodium, Na, 11 alkali metal1, 3, s22.98976928(2) g·mol−1 [Ne] 3s1 2,8,1 (Image) solid 0.968 g·cm−3 0.927 g·cm−3 370.87 K,97.72 °C,207.9 °F 1156 K,883 °C,1621 °F (extrapolated) 2573 K, 35 MPa 2.60 kJ·mol−1 97.42 kJ·mol−1 (25 °C) 28.230 J·mol−1·K−1+1, -1 (strongly basic oxide) 0.93 (Pauling scale) 1st: 495.8 kJ·mol−1186 pm166±9 pm 227 pm body-centered cubic paramagnetic (20 °C) 47.7 nΩ·m (300 K) 142 W·m−1·K−1 (25 °C) 71 µm·m−1·K−1 (20 °C) 3200 m/s 10 GPa 3.3 GPa 6.3 GPa 0.5 0.69 MPa 7440-23-5 Sodium (pronounced /ˈsoʊdiəm/) is a metallic element with a symbol Na (from Latin natrium or Arabic natrun) and atomic number 11. It is a soft, silvery-white, highly reactive metal and is a member of the alkali metals within "group 1" (formerly known as ‘group IA’). It has only one stable isotope, 23Na. Elemental sodium was first isolated by Sir Humphry Davy in 1806 by passing an electric current through molten sodium hydroxide. Elemental sodium does not occur naturally on Earth, but quickly oxidizes in air and is violently reactive with water, so it must be stored in an inert medium, such as a liquid hydrocarbon. The free metal is used for some chemical synthesis and heat transfer applications. Sodium ion is soluble in water in nearly all of its compounds, and is thus present in great quantities in the Earth's oceans and other stagnant bodies of water. In these bodies it is mostly counterbalanced by the chloride ion, causing evaporated ocean water solids to consist mostly of sodium chloride, or common table salt. Sodium ion is also a component of many minerals. Sodium is an essential element for all animal life and for some plant species. In animals, sodium ions are used in opposition to potassium ions, to allow the organism to build up an electrostatic charge on cell membranes, and thus allow transmission of nerve impulses when the charge is allowed to dissipate by a moving wave of voltage change. Sodium is thus classified as a “dietary inorganic macro-mineral” for animals. Sodium's relative rarity on land is due to its solubility in water, thus causing it to be leached into bodies of long-standing water by rainfall. Such is its relatively large requirement in animals, in contrast to its relative scarcity in many inland soils, that herbivorous land animals have developed a special taste receptor for sodium ion.
Sodium
3
Characteristics At room temperature, sodium metal is soft enough that it can be cut with a knife. In air, the bright silvery luster of freshly exposed sodium will rapidly tarnish. The density of alkali metals generally increases with increasing atomic number, but sodium is denser than potassium.
Chemical properties Compared with other alkali metals, sodium is generally less reactive than potassium and more reactive than lithium,[2] in accordance with "periodic law": for example, their reaction in water, chlorine gas, etc.; Sodium reacts exothermically with water: small pea-sized pieces will bounce across the surface of the water until they are consumed by it, whereas large pieces will explode. While sodium reacts with water at room temperature, the sodium piece melts with the heat of the reaction to form a sphere, if the reacting sodium piece is large enough. The reaction with water produces very caustic sodium hydroxide (lye) and highly flammable hydrogen gas. These are extreme hazards (see Precautions section below). When burned in air, sodium forms sodium peroxide Na2O2, or with limited oxygen, the oxide Na2O (unlike lithium, the nitride is not formed). If burned in oxygen under pressure, sodium superoxide NaO2 will be produced. In chemistry, most sodium compounds are considered soluble but nature Sodium metal provides examples of many insoluble sodium compounds such as the (approx 10g) under feldspars. There are other insoluble sodium salts such as sodium oil bismuthate NaBiO3, sodium octamolybdate Na2Mo8O25• 4H2O, sodium thioplatinate Na4Pt3S6, sodium uranate Na2UO4. Sodium meta-antimonate's 2NaSbO3•7H2O solubility is 0.3g/L as is the pyro form Na2H2Sb2O7• H2O of this salt. Sodium metaphosphate NaPO3 has a soluble and an insoluble form.[3]
Compounds Sodium compounds are important to the chemical, glass, metal, paper, petroleum, soap, and textile industries. Hard soaps are generally sodium salt of certain fatty acids (potassium produces softer or liquid soaps).[4] The sodium compounds that are the most important to industries are common salt (NaCl), soda ash (Na2CO3), baking soda (NaHCO3), caustic soda (NaOH), sodium nitrate (NaNO3), di- and tri-sodium phosphates, sodium thiosulfate (hypo, Na2S2O3 · 5H2O), and borax (Na2B4O7 · 10H2O).[4]
Spectroscopy When sodium or its compounds are introduced into a flame, they turn the flame a bright yellow color. Sodium spectral lines.
Sodium
4
A low pressure sodium/sodium oxide (LPS/SOX) streetlamp at full power (detail)
One notable atomic spectral line of sodium vapor is the so-called D-line, which may be observed directly as the sodium flame-test line (see Applications) and also the major light output of low-pressure sodium lamps (these produce an unnatural yellow, rather than the peach-colored glow of high pressure lamps). The D-line is one of the classified Fraunhofer lines observed in the visible spectrum of the Sun's electromagnetic radiation. Sodium vapor in the upper layers of the Sun creates a dark line in the emitted spectrum of electromagnetic radiation by absorbing visible light in a band of wavelengths around 589.5 nm. This wavelength corresponds to transitions in atomic sodium in which the valence-electron transitions from a 3p to 3s electronic state. Closer examination of the visible spectrum of atomic sodium reveals that the D-line actually consists of two lines called the D1 and D2 lines at 589.6 nm and 589.0 nm, respectively. This fine structure results from a spin-orbit interaction of the valence electron in the 3p electronic state. The spin-orbit interaction couples the spin angular momentum and orbital angular momentum of a 3p electron to form two states that are respectively notated
as
and
in
the
LS
coupling scheme. The 3s state of the electron gives rise to a single state which is notated as
in
the LS coupling scheme. The D1-line results from an electronic transition between and
lower state
upper state. The D2-line results from
an electronic transition between state
and
upper
state.
lower Even
closer
examination of the visible spectrum of atomic sodium would reveal that the D-line actually consists A FASOR tuned to the D2A component of the sodium D line, used at the Starfire Optical Range to excite sodium atoms in the upper atmosphere.
hyperfine levels.[5]
of a lot more than two lines. These lines are associated with hyperfine structure of the 3p upper states and 3s lower states. Many different transitions involving visible light near 589.5 nm may occur between the different upper and lower
[6]
A practical use for lasers which work at the sodium D-line transition (see FASOR illustration) is to create artificial laser guide stars (artificial star-like images from sodium in the upper atmosphere) which assist in the adaptive optics for large land-based visible light telescopes.
Sodium
5
Isotopes Thirteen isotopes of sodium have been recognized. The only stable isotope is 23Na. Sodium has two radioactive cosmogenic isotopes which are also the two isotopes with longest half life, 22Na, with a half-life of 2.6 years and 24Na with a half-life of 15 hours. All other isotopes have a half life of less than one minute.[7] Acute neutron radiation exposure (e.g., from a nuclear criticality accident) converts some of the stable 23Na in human blood plasma to 24Na. By measuring the concentration of this isotope, the neutron radiation dosage to the victim can be computed.[8]
History Salt has been an important commodity in human activities, as testified by the English word salary, referring to salarium, the wafers of salt sometimes given to Roman soldiers along with their other wages. In medieval Europe a compound of sodium with the Latin name of sodanum was used as a headache remedy. The name sodium probably originates from the Arabic word suda meaning headache as the headache-alleviating properties of sodium carbonate or soda were well known in early times.[9] Sodium's
The flame test for sodium displays a brilliantly bright yellow emission due to the so called "sodium D-lines" at 588.9950 and 589.5924 nanometers.
chemical
abbreviation
Na
was
first
published by Jöns Jakob Berzelius in his system of atomic symbols (Thomas Thomson, Annals of Philosophy[10] ) and is a contraction of the element's new Latin name natrium which refers to the Egyptian natron,[11] the word for a natural mineral salt whose primary ingredient is hydrated sodium carbonate. Hydrated sodium carbonate historically had several important industrial and household uses later eclipsed by soda ash, baking soda and other sodium compounds.
Although sodium (sometimes called "soda" in English) has long been recognized in compounds, it was not isolated until 1807 by Sir Humphry Davy through the electrolysis of caustic soda.[12] Sodium imparts an intense yellow color to flames. As early as 1860, Kirchhoff and Bunsen noted the high sensitivity that a flame test for sodium could give. They state in Annalen der Physik und der Chemie in the paper "Chemical Analysis by Observation of Spectra": In a corner of our 60 cu.m. room farthest away from the apparatus, we exploded 3 mg. of sodium chlorate with milk sugar while observing the nonluminous flame before the slit. After a while, it glowed a bright yellow and showed a strong sodium line that disappeared only after 10 minutes. From the weight of the sodium salt and the volume of air in the room, we easily calculate that one part by weight of air could not contain more than 1/20 millionth
Sodium
6
weight of sodium.
Occurrence Owing to its high reactivity, sodium is found in nature only as a compound and never as the free element. Sodium makes up about 2.6% by weight of the Earth's crust, making it the sixth most abundant element overall[13] and the most abundant alkali metal. Sodium is found in many different minerals, of which the most common is ordinary salt (sodium chloride), which occurs in vast quantities dissolved in seawater, as well as in solid deposits (halite). Others include amphibole, cryolite, soda niter and zeolite. Sodium is relatively abundant in stars and the D spectral lines of this element are among the most prominent in star light. Though elemental sodium has a rather high vaporization temperature, its relatively high abundance and very intense spectral lines have allowed its presence to be detected by ground telescopes and confirmed by spacecraft (Mariner 10 and MESSENGER) in the thin atmosphere of the planet Mercury.[14]
Commercial production Sodium was first produced commercially in 1855 by thermal reduction of sodium carbonate with carbon at 1100 °C, in what is known as the Deville process.[15] Na2CO3 (liquid) + 2 C (solid) → 2 Na (vapor) + 3 CO (gas). A process based on the reduction of sodium hydroxide was developed in 1886.[15] Sodium is now produced commercially through the electrolysis of liquid sodium chloride, based on a process patented in 1924.[16] [17] This is done in a Downs Cell in which the NaCl is mixed with calcium chloride to lower the melting point below 700 °C. As calcium is less electropositive than sodium, no calcium will be formed at the anode. This method is less expensive than the previous Castner process of electrolyzing sodium hydroxide. Very pure sodium can be isolated by the thermal decomposition of sodium azide.[18] Metallic sodium costs about 15 to 20 US cents per pound (US$0.30/kg to US$0.45/kg) in 1997, but reagent grade (ACS) sodium cost about US$35 per pound (US$75/kg) in 1990.
Applications Sodium in its metallic form can be used to refine some reactive metals, such as zirconium and potassium, from their compounds. This alkali metal as the Na+ ion is vital to animal life. Other uses: • In certain alloys to improve their structure. • In soap, in combination with fatty acids. Sodium soaps are harder (higher melting) soaps than potassium soaps. • To descale metal (make its surface smooth). • To purify molten metals. • In some medicine formulations, the salt form of the active ingredient usually with sodium or potassium is a common modification to improve bioavailability. • In sodium vapor lamps, an efficient means of producing light from electricity (see the picture), often used for street lighting in cities. Low-pressure sodium lamps give a distinctive yellow-orange light which consists primarily of the twin sodium D lines. High-pressure sodium lamps give a more natural peach-colored light, composed of
Sodium wavelengths spread much more widely across the spectrum. • As a heat transfer fluid in some types of nuclear reactors and inside the hollow valves of high-performance internal combustion engines. • Sodium chloride (NaCl), a compound of sodium ions and chloride ions, is an important heat transfer material. • In organic synthesis, sodium is used as a reducing agent, for example in the Birch reduction. • In chemistry, sodium is often used either alone or with potassium in an alloy, NaK as a desiccant for drying solvents. Used with benzophenone, it forms an intense blue coloration when the solvent is dry and oxygen-free.
Biological role In maintenance of body fluid volume in animals The serum sodium and urine sodium play important roles in medicine, both in the maintenance of sodium and total body fluid homeostasis, and in the diagnosis of disorders causing homeostatic disruption of salt/sodium and water balance. In mammals, decreases in blood pressure and decreases in sodium concentration sensed within the kidney result in the production of renin, a hormone which acts in a number of ways, one of them being to act indirectly to cause the generation of aldosterone, a hormone which decreases the excretion of sodium in the urine. As the body of the mammal retains more sodium, other osmoregulation systems which sense osmotic pressure in part from the concentration of sodium and water in the blood, act to generate antidiuretic hormone. This, in turn, which causes the body to retain water, thus helping to restoring the body's total amount of fluid. There is also a counterbalancing system, which senses volume. As fluid is retained, receptors in the heart and vessels which sense distension and pressure, cause production of atrial natriuretic peptide, which is named in part for the Latin word for sodium. This hormone acts in various ways to cause the body to lose sodium in the urine. This causes the body's osmotic balance to drop (as low concentration of sodium is sensed directly), which in turn causes the osmoregulation system to excrete the "excess" water. The net effect is to return the body's total fluid levels back toward normal.
In maintenance of resting electrical potential in excitable tissues in animals Sodium cations are important in neuron (brain and nerve) function, and in influencing osmotic balance between cells and the interstitial fluid, with their distribution mediated in all animals (but not in all plants) by the so-called Na+/K+-ATPase pump.[19] Sodium is the chief cation in fluid residing outside cells in the mammalian body (the so-called extracellular compartment), with relatively little sodium residing inside cells. The volume of extracellular fluid is typically 15 litres in a 70 kg human, and the 50 grams of sodium it contains is about 90% of the body's total sodium content.
7
Sodium
Dietary uses The most common sodium salt, sodium chloride (table salt), is used for seasoning and warm-climate food preservation, such as pickling and making jerky (the high osmotic content of salt inhibits bacterial and fungal growth). The human requirement for sodium in the diet is about 500 mg per day,[20] which is typically less than a tenth as much as many diets "seasoned to taste." Most people consume far more sodium than is physiologically needed. For certain people with salt-sensitive blood pressure, this extra intake may cause a harmful effect on health. However, low sodium intake may lead to sodium deficiency.
Precautions Extreme care is required in handling elemental/metallic sodium. Sodium is potentially explosive in water (depending on quantity) and is a corrosive substance, since it is rapidly converted to sodium hydroxide on contact with moisture. The powdered form may combust spontaneously in air or oxygen. Sodium must be stored either in an inert (oxygen and moisture free) atmosphere (such as nitrogen or argon), or under a liquid hydrocarbon such as mineral oil or kerosene. The reaction of sodium and water is a familiar one in chemistry labs, and is reasonably safe if amounts of sodium smaller than a pencil eraser are used and the reaction is done behind a plastic shield by people wearing eye protection. However, the sodium-water reaction does not scale up well, and is treacherous when larger amounts of sodium are used. Larger pieces of sodium melt under the heat of the reaction, and the molten ball of metal is buoyed up by hydrogen and may appear to be stably reacting with water, until splashing covers more of the reaction mass, causing thermal runaway and an explosion which scatters molten sodium, lye solution, and sometimes flame. (18.5 g explosion [21]) This behavior is unpredictable, and among the alkali metals it is usually sodium which invites this surprise phenomenon, because lithium is not reactive enough to do it, and potassium is so reactive that chemistry students are not tempted to try the reaction with larger potassium pieces. Sodium is much more reactive than magnesium; a reactivity which can be further enhanced due to sodium's much lower melting point. When sodium catches fire in air (as opposed to just the hydrogen gas generated from water by means of its reaction with sodium) it more easily produces temperatures high enough to melt the sodium, exposing more of its surface to the air and spreading the fire. Few common fire extinguishers work on sodium fires. Water, of course, exacerbates sodium fires, as do water-based foams. CO2 and Halon are often ineffective on sodium fires, which reignite when the extinguisher dissipates. Among the very few materials effective on a sodium fire are Pyromet and Met-L-X. Pyromet is a NaCl/(NH4)2HPO4 mix, with flow/anti-clump agents. It smothers the fire, drains away heat, and melts to form an impermeable crust. This is the standard dry-powder canister fire extinguisher for all classes of fires. Met-L-X is mostly sodium chloride, NaCl, with approximately 5% Saran plastic as a crust-former, and flow/anti-clumping agents. It is most commonly hand-applied, with a scoop. Other extreme fire extinguishing materials include Lith+, a graphite based dry powder with an organophosphate flame retardant; and Na+, a Na2CO3-based material. Because of the reaction scale problems discussed above, disposing of large quantities of sodium (more than 10 to 100 grams) must be done through a licensed hazardous materials disposer. Smaller quantities may be broken up and neutralized carefully with ethanol
8
Sodium
9
(which has a much slower reaction than water), or even methanol (where the reaction is more rapid than ethanol's but still less than in water), but care should nevertheless be taken, as the caustic products from the ethanol or methanol reaction are just as hazardous to eyes and skin as those from water. After the alcohol reaction appears complete, and all pieces of reaction debris have been broken up or dissolved, a mixture of alcohol and water, then pure water, may then be carefully used for a final cleaning. This should be allowed to stand a few minutes until the reaction products are diluted more thoroughly and flushed down the drain. The purpose of the final water soaking and washing of any reaction mass or container which may contain sodium, is to ensure that alcohol does not carry unreacted sodium into the sink trap, where a water reaction may generate hydrogen in the trap space which can then be potentially ignited, causing a confined sink trap explosion.
See also • Alkali metals • Sodium compounds
External links • • • • • •
The Periodic Table of Videos - Sodium [22] Etymology of "natrium" - source of symbol Na [23] WebElements.com – Sodium [24] The Wooden Periodic Table Table's Entry on Sodium [25] Dietary Sodium [26] Sodium isotopes data from The Berkeley Laboratory Isotopes Project's
[27]
References [1] Endt, P. M. ENDT, ,1 (1990) (12/1990). "Energy levels of A = 21-44 nuclei (VII)". Nuclear Physics A 521: 1. doi: 10.1016/0375-9474(90)90598-G (http:/ / dx. doi. org/ 10. 1016/ 0375-9474(90)90598-G). [2] Prof. N. De Leon. " Reactivity of Alkali Metals (http:/ / www. iun. edu/ ~cpanhd/ C101webnotes/ modern-atomic-theory/ alkali-reac. html)". Indiana University Northwest. . Retrieved 2007-12-07. [3] Lange's Handbook of Chemistry [4] Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils; (1985). "Natrium" (in German). Lehrbuch der Anorganischen Chemie (91–100 ed.). Walter de Gruyter. pp. 931–943. ISBN 3-11-007511-3. [5] Citron, M. L., et al. (1977). " Experimental study of power broadening in a two level atom (http:/ / prola. aps. org/ abstract/ PRA/ v16/ i4/ p1507_1)". Physical Review A 16: 1507. doi: 10.1103/PhysRevA.16.1507 (http:/ / dx. doi. org/ 10. 1103/ PhysRevA. 16. 1507). . [6] Daniel A. Steck. " Sodium D. Line Data (http:/ / george. ph. utexas. edu/ ~dsteck/ alkalidata/ sodiumnumbers. pdf)" (PDF). Los Alamos National Laboratory (technical report). . [7] Audi, Georges (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3–128. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [8] Sanders, F. W.; Auxier, J. A. (1962). " Neutron Activation of Sodium in Anthropomorphous Phantoms (http:/ / www. health-physics. com/ pt/ re/ healthphys/ abstract. 00004032-196208000-00005. htm)". Health Physics 8 (4): 371–379. doi: 10.1097/00004032-196208000-00005 (http:/ / dx. doi. org/ 10. 1097/ 00004032-196208000-00005). . [9] David E. Newton, Chemical Elements, ISBN 0-7876-2847-6 [10] van der Krogt, Peter. " Elementymology & Elements Multidict (http:/ / www. vanderkrogt. net/ elements/ elem/ na. html)". . Retrieved 2007-06-08. [11] Newton [12] Davy, Humphry (1808). " On some new Phenomena of Chemical Changes produced by Electricity, particularly the Decomposition of the fixed Alkalies, and the Exhibition of the new Substances, which constitute their Bases (http:/ / books. google. com/ books?id=Kg9GAAAAMAAJ)". Philosophical Transactions of the Royal Society of
Sodium
10
London 98: 1–45. doi: 10.1098/rstl.1808.0001 (http:/ / dx. doi. org/ 10. 1098/ rstl. 1808. 0001). . [13] CRC Handbook of Chemistry and Physics, 2004 [14] " Sodium found in Mercury's atmosphere (http:/ / findarticles. com/ p/ articles/ mi_m1200/ is_v128/ ai_3898126)". BNET. 1985-08-17. . Retrieved 2008-09-18. [15] Eggeman, Tim. Sodium and Sodium Alloys. Kirk-Othmer Encyclopedia of Chemical Technology. John Wiley & Sons, Inc. Published online 2007. doi: 10.1002/0471238961.1915040912051311.a01.pub2 (http:/ / dx. doi. org/ 10. 1002/ 0471238961. 1915040912051311. a01. pub2) [16] Pauling, Linus, General Chemistry, 1970 ed., Dover Publications [17] " Los Alamos National Laboratory – Sodium (http:/ / periodic. lanl. gov/ elements/ 11. html)". . Retrieved 2007-06-08. [18] Merck Index, 9th ed., monograph 8325 [19] Campbell, Neil (1987). Biology. Menlo Park, Calif.: Benjamin/Cummings Pub. Co.. pp. 795. ISBN 0-8053-1840-2. [20] Implementing recommendations for dietary salt reduction: Where are we?. DIANE Publishing. ISBN 1428929096. [21] [22] [23] [24] [25] [26] [27]
http:/ / video. google. de/ videoplay?docid=-2158222101210607510& q=sodium http:/ / www. youtube. com/ watch?v=YvSkXd_VVYk http:/ / www. balashon. com/ 2008/ 07/ neter-and-nitrogen. html http:/ / www. webelements. com/ sodium/ http:/ / www. theodoregray. com/ PeriodicTable/ Elements/ 011/ index. html http:/ / www. americanheart. org/ presenter. jhtml?identifier=4708 http:/ / ie. lbl. gov/ education/ parent/ Na_iso. htm
Article Sources and Contributors
Article Sources and Contributors Sodium Source: http://en.wikipedia.org/w/index.php?oldid=306291367 Contributors: (jarbarf), -Ozone-, 1Shaggy1, 21655, 2D, 454Casull, AAA!, ABF, Abrutt, Academic Challenger, Acroterion, AdamRetchless, Adashiel, Addshore, Adrian.benko, Aeluwas, Agateller, Agüeybaná, Ahoerstemeier, Aircavbob, Alansohn, Aleenf1, Alexfusco5, Alexwcovington, Alison, All Is One, Allstarecho, Alton, Ameliorate!, AndreasJS, Andres, AndrewMcQ, Andrewa, Andrewrost3241981, AnnaFrance, Anonymous101, Antandrus, Antonio Lopez, Anwar saadat, Aoganov, Arcadian, Archimerged, Arwack, Arx Fortis, Asb002, Ascend, Ascidian, Ashleybabe92, AtomicDragon, Atrzcins, Balwan, Bchalfin, Bcorr, Beetstra, Bella Swan, BenFrantzDale, Benbest, Bensaccount, BernardM, Bihco, Binary TSO, Blue520, BlueEarth, Bobo The Ninja, Bobo192, Bomac, Bongwarrior, Bookandcoffee, Bookwyrm404, Borbrav, Borislav Dopudja, Bornhj, Brian Crawford, Bryan Derksen, Bubba hotep, Bubblecourt, Bucketsofg, C.lettinga, CHawke, CYD, Can't sleep, clown will eat me, CanisRufus, Canyoyoman, Capricorn42, Captain panda, Carlj7, Carnildo, Cbh, Cburnett, Celarnor, Cfrehr, Chairman S., Chaser, ChemGardener, Chris the speller, Chriswaterguy, Chun-hian, Chunky Rice, Clemwang, Click23, Closedmouth, Cloud Strife, Colbuckshot, Cometstyles, Conversion script, Coppertwig, Corpx, Cow Revolution, Cowman109, Cremepuff222, Crownmethod, Cureden, D, DB, DMacks, Daniel 1992, Daniel Olsen, DanielCD, Darguz Parsilvan, DarrenHarold, Darrien, Daverocks, David Latapie, Dbtfz, DeadEyeArrow, Deglr6328, Delta G, Deor, DerHexer, Digger318, Digitalme, Dina, Discospinster, Dlohcierekim, Docboat, Donarreiskoffer, DougsTech, Doulos Christos, Dppowell, DrBob, Dracosummoner, DragonflySixtyseven, Dreadstar, Drini, Droll, Drunkenmonkey, Dysepsion, EJF, ESkog, Eddideigel, Edgar181, Eenu, El C, Eldin raigmore, Elektron, Ellywa, Emperorbma, Encyclopedia77, Enigmaman, Enviroboy, Eob, Epbr123, Epolk, Eric119, Erik Zachte, Erik9, Eternal Pink, Euchiasmus, Everyking, Evil Monkey, Excirial, Faradayplank, Farosdaughter, Fdp, Femto, Fieldday-sunday, Firzen67, Flo422, Fonzy, FourteenDays, FrancoGG, Friginator, Func, Fvw, GHe, Gail, Galoubet, Gdo01, Gene Nygaard, Giftlite, Ginkgo100, Glen, Glenn, Gman124, Gmmaple, Gogo Dodo, GraYoshi2x, Gravecat, GreatMizuti, Greenhorn1, Grendelkhan, Greylion, Guitarmankev1, Gurch, Gwen Gale, Gwernol, Hadal, Hak-kâ-ngìn, Hankwang, Hatmatbbat10, Hda3ku, HenryLi, Herbee, HexaChord, Hodiedodie, Hppybob, Hqb, Hraefen, Hu, Husayt, Hydrogen Iodide, I already forgot, II MusLiM HyBRiD II, Iamsam478, Icairns, Iloveorange, Insanity Incarnate, Instinct, Invincible Ninja, Iridescent, Iris lorain, Island, Itub, Ixfd64, J.delanoy, JCO312, JForget, Jagged 85, Jamesontai, JamieS93, Jaraalbe, Jauerback, JavierMC, Jcw69, Jdurg, Jeff G., Jeffreyapm, JericoReyes, Jhbayless, JimQ, Jj137, JoanneB, John, John David Wright, John254, Jonnyl94, Jose77, Jpeeling, JuJube, Juliancolton, Jumbuck, JustShin, Kaisershatner, Karenjc, Karl-Henner, Karlhahn, Karn, Katalaveno, Kazrak, Keandobus, Kesac, Kevin Dufendach, Kevindavid, Kfrohlinger, Killiondude, Klausness, Knowledge Seeker, KnowledgeOfSelf, Knutux, KonradG, Koobalocks, Kpjas, Kralizec!, Ktsquare, Kungfuadam, Kuru, Kurykh, Kwamikagami, Kwsn, LaMenta3, Larry laptop, LarryMorseDCOhio, Laudaka, Lindmere, LittleOldMe, Lizadances, Lone star06, Loren.wilton, LorenzoB, Lradrama, LuigiManiac, Luvy, MER-C, MFago, MZMcBride, Mac Davis, Malcolm Farmer, Man vyi, Manofradio, Marek69, Marnanel, Martian.knight, Materialscientist, Mav, Maxberners, May0208, Mayamaya305, Meekywiki, Menchi, Metal Militia, Mexcellent, Mgimpel, Mgonz5, Mhking, Midgley, Mikael Häggström, Mike133742, Mike2vil, Mikemastelero, Minesweeper, MisterWiki, Mithunc, Mjp797, Mnuskey, Mononomic, Moreschi, Mr.Z-man, Mwanner, Mxn, MysteryDog, Nakon, Nateinbliss, NathanHurst, NawlinWiki, Nehrams2020, Neillawrence, Neurolysis, NewEnglandYankee, Nick, Nick UA, Nickgully, Nihiltres, Ninja Wizard, Nivix, Nsaa, NuclearWarfare, Num3boy, Ocatecir, Oda Mari, Ohnoitsjamie, Oliver Lineham, Onceonthisisland, Oore, Ossmann, Owen, Oxymoron83, Oysterguitarist, Padfoot addict, Pairadox, ParisianBlade, Patstuart, PaulHanson, Pedriana, Peruvianllama, Peter bertok, Pharos, Phil Boswell, Philip Trueman, Phillman5, Piano non troppo, Pilotguy, Pkl$$$123, Plantsurfer, PoliteCarbide, Poolkris, Postdlf, Psyche825, Puchiko, Quadell, Quantockgoblin, Quintote, Quinwound, RDBrown, RJaguar3, RSido, RTC, Ranveig, Raymondwinn, Rchandra, Rdsmith4, Recognizance, Red Thunder, Reddi, Remember, Res2216firestar, Rettetast, RexNL, Riana, Rich Farmbrough, Richard D. 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11
Magnesium
1
Magnesium sodium ← magnesium → aluminiumBe ↑ Mg ↓ Ca
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
12Mg Periodic table
Appearance silvery white solid
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
701
773
861
971
1132
1361
Magnesium
2
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1450.7 kJ·mol−1 3rd: 7732.7 kJ·mol−1Atomic radiusCovalent radiusVan der Waals radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusPoisson ratioMohs hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of magnesium iso
N.A.
half-life
24
78.99%
24
25
10%
25
26
11.01%
26
Mg Mg Mg
DM
DE (MeV)
DP
Mg is stable with 12 neutron Mg is stable with 13 neutron Mg is stable with 14 neutron
magnesium, Mg, 12 alkaline earth metal2, 3, s24.3050(6) g·mol−1 [Ne] 3s2 2, 8, 2 (Image) solid 1.738 g·cm−3 1.584 g·cm−3 923 K,650 °C,1202 °F 1363 K,1091 °C,1994 °F 8.48 kJ·mol−1 128 kJ·mol−1 (25 °C) 24.869 J·mol−1·K−12, 1 [1] (strongly basic oxide) 1.31 (Pauling scale) 1st: 737.7 kJ·mol−1160 pm141±7 pm 173 pm hexagonal paramagnetic (20 °C) 43.9 nΩ·m (300 K) 156 W·m−1·K−1 (25 °C) 24.8 µm·m−1·K−1 (r.t.) (annealed) 4940 m·s−1 45 GPa 17 GPa 45 GPa 0.290 2.5 260 MPa 7439-95-4 Magnesium (pronounced /mæɡˈniːziəm/) is a chemical element with the symbol Mg, atomic number 12, atomic weight 24.3050 and common oxidation number +2. Magnesium, an alkaline earth metal, is the ninth most abundant element in the universe by mass.[2] This preponderance of magnesium is related to the fact that it is easily built up in supernova stars from a sequential addition of three helium nuclei to carbon (which in turn is made from a single reaction between three helium nuclei at once). Magnesium constitutes about 2% of the Earth's crust by mass, which makes it the eighth most abundant element in the crust.[3] Magnesium ion's high solubility in water helps ensure that it is the third most abundant element dissolved in seawater.[4] Magnesium is the 11th most abundant element by mass in the human body; its ions are essential to all living cells, where they play a major role in manipulating important biological polyphosphate compounds like ATP, DNA, and RNA. Hundreds of enzymes thus require magnesium ions in order to function. Magnesium is also the metallic ion at the center of chlorophyll, and is thus a common additive to fertilizers.[5] Magnesium compounds are used medicinally as common laxatives, antacids (i.e., milk of magnesia), and in a number of situations where stabilization of abnormal nerve excitation and blood vessel spasm is required (i.e., to treat eclampsia). Magnesium ions are sour to the taste, and in low concentrations help to impart a natural tartness to fresh mineral waters. The free element (metal) is not found naturally on Earth, since it is highly reactive (though once produced, is coated in a thin layer of oxide—see passivation—which partly masks this reactivity). The free metal burns with a characteristic brilliant white light, making it a useful ingredient in flares. The metal is now mainly obtained by electrolysis of magnesium salts obtained from brine. Commercially, the chief use for the metal is as an alloying agent to make aluminium-magnesium alloys, sometimes called "magnalium" or "magnelium". Since magnesium is less dense than aluminium, these alloys are prized for their relative lightness and strength.
Magnesium
Notable characteristics Elemental magnesium is a fairly strong, silvery-white, light-weight metal (two thirds the density of aluminium). It tarnishes slightly when exposed to air, although unlike the alkaline metals, storage in an oxygen-free environment is unnecessary because magnesium is protected by a thin layer of oxide which is fairly impermeable and hard to remove. Like its lower periodic table group neighbor calcium, magnesium reacts with water at room temperature, though it reacts much more slowly than calcium. When it is submerged in water, hydrogen bubbles will almost unnoticeably begin to form on the surface of the metal, though if powdered it will react much more rapidly. The reaction will occur faster with higher temperatures (see precautions). Magnesium also reacts exothermically with most acids, such as hydrochloric acid (HCl). As with aluminium, zinc and many other metals, the reaction with hydrochloric acid produces the chloride of the metal and releases hydrogen gas. Magnesium is a highly flammable metal, but while it is easy to ignite when powdered or shaved into thin strips, it is difficult to ignite in mass or bulk. Once ignited, it is difficult to extinguish, being able to burn in both nitrogen (forming magnesium nitride), and carbon dioxide (forming magnesium oxide and carbon). On burning in air, magnesium produces a brilliant white light. Thus magnesium powder (flash powder) was used as a source of illumination in the early days of photography. Later, magnesium ribbon was used in electrically ignited flash bulbs. Magnesium powder is used in the manufacture of fireworks and marine flares where a brilliant white light is required. Flame temperatures of magnesium and magnesium alloys can reach 1371 °C (2500 °F), although flame height above the burning metal is usually less than 300 mm (12 in).[6] Magnesium compounds are typically white crystals. Most are soluble in water, providing the sour-tasting magnesium ion Mg2+. Small amounts of dissolved magnesium ion contributes to the tartness and taste of natural waters. Magnesium ion in large amounts is an ionic laxative, and magnesium sulfate (Epsom salts) is sometimes used for this purpose. So-called "milk of magnesia" is a water suspension of one of the few insoluble magnesium compounds, magnesium hydroxide. The undissolved particles give rise to its appearance and name. Milk of magnesia is a mild base commonly used as an antacid.
Applications As the metal
3
Magnesium
4 Magnesium is the third most commonly used structural metal, following steel and aluminium.
An unusual application of magnesium as an illumination source while wakeskating in 1931
Magnesium compounds, primarily magnesium oxide (MgO), are used mainly as refractory material in furnace linings for producing iron, steel, nonferrous metals, glass and cement. Magnesium oxide and other compounds also are used in agricultural, chemical and construction industries. As a metal, this element's principal use is as an alloying additive to aluminium with these aluminium-magnesium alloys being used mainly for beverage cans. Magnesium, in its purest form, can be compared with
aluminium, and is strong and light, so it is used in several high volume part manufacturing applications, including automotive and truck components. Specialty, high-grade car wheels of magnesium alloy are called "mag wheels". In 1957 a Corvette SS, designed for racing, was constructed with magnesium body panels. An earlier Mercedes-Benz race car model, the Mercedes-Benz 300 SLR, had a body made from Elektron, a magnesium alloy; these cars ran (with successes) at Le Mans, the Mille Miglia, and other world-class race events in 1955 (though one was involved in the single worst accident in auto racing history, in terms of human casualties, at the Le Mans race.) Porsche's all-out quest to decrease the weight of their race cars led to the use of magnesium frames in the famous 917/053 which won Le Mans in 1971, and still holds the absolute distance record. The 917/30 Can-Am car also featured a magnesium spaceframe, helping it to make the most of its prodigious 1100-1500hp. Volkswagen has used magnesium in its engine components for many years. For a long time, Porsche used magnesium alloy for its engine blocks due to the weight advantage. There is renewed interest in magnesium engine blocks, as featured in the 2006 BMW 325i and 330i models. The BMW engine uses an aluminium alloy insert for the cylinder walls and cooling jackets surrounded by a high temperature magnesium alloy AJ62A. The application of magnesium AE44 alloy in the 2006 Corvette Z06 engine cradle has advanced the technology of designing robust automotive parts in magnesium. Both of these alloys are recent developments in high temperature low creep magnesium alloys. Mitsubishi Motors also uses magnesium (branded magnesium alloy) for its paddle shifters. The general strategy for such alloys is to form intermetallic precipitates at the grain boundaries, for example by adding mischmetal or calcium.[7] New alloy development and lower costs, which are becoming competitive to aluminium, will further the number of automotive applications.
Magnesium
5 The second application field of magnesium is electronic devices. Due to low weight, good mechanical and electrical properties, magnesium is widely used for manufacturing of mobile phones, laptop computers, cameras, and other electronic components.
Historically, magnesium was one of the main aerospace construction metals and was used for German military aircraft as early as World War I and extensively for German aircraft in World War II. The Germans coined Products made of magnesium: the name 'Elektron' for magnesium alloy which is still firestarter and shavings, sharpener, used today. Due to perceived hazards with magnesium magnesium ribbon parts in the event of fire, the application of magnesium in the commercial aerospace industry was generally restricted to engine related components. Currently the use of magnesium alloys in aerospace is increasing, mostly driven by the increasing importance of fuel economy and the need to reduce weight. The development and testing of new magnesium alloys continues, notably Elektron 21 which has successfully undergone extensive aerospace testing for suitability in engine, internal and airframe components. The European Community runs three R&D magnesium projects in the Aerospace priority of Six Framework Program. • Incendiary use: Magnesium is flammable, burning at a temperature of approximately 2500 K (2200 °C, 4000 °F), and the autoignition temperature of magnesium is approximately 744 K (473 °C, 883 °F) in air. The extremely high temperature at which magnesium burns makes it a handy tool for starting emergency fires during outdoor recreation. Other related uses include flashlight photography, flares, pyrotechnics, fireworks sparklers, and incendiary bombs. Magnesium is also used: • To remove sulfur from iron and steel. • To refine titanium in the Kroll process.
Magnesium firestarter (in left hand), used with a pocket knife and flint to create sparks which ignite the shavings
• To photoengrave plates in the printing industry. • To combine in alloys, where this metal is essential for airplane and missile construction. • In the form of turnings or ribbons, to prepare Grignard reagents, which are useful in organic synthesis. • As an alloying agent, improving the mechanical, fabrication and welding characteristics of aluminium. • As an additive agent in conventional propellants and the production of nodular graphite in cast iron. • As a reducing agent for the production of uranium and other metals from their salts. • As a desiccant, since it easily reacts with water. • As a sacrificial (galvanic) anode to protect underground tanks, pipelines, buried structures, and water heaters.
Magnesium
In magnesium compounds • The magnesium ion is necessary for all life (see magnesium in biology), so magnesium salts are an additive for foods, fertilizers (Mg is a component of chlorophyll), and culture media. • Magnesium hydroxide is used in milk of magnesia, its chloride, oxide, gluconate, malate, orotate and citrate used as oral magnesium supplements, and its sulfate (Epsom salts) for various purposes in medicine, and elsewhere (see the article for more). Oral magnesium supplements have been claimed to be therapeutic for some individuals who suffer from Restless Leg Syndrome (RLS). • Magnesium borate, magnesium salicylate and magnesium sulfate are used as antiseptics. • Magnesium bromide is used as a mild sedative (this action is due to the bromide, not the magnesium). • Dead-burned magnesite is used for refractory purposes such as brick and liners in furnaces and converters. • Magnesium carbonate (MgCO3) powder is also used by athletes, such as gymnasts and weightlifters, to improve the grip on objects – the apparatus or lifting bar. • Magnesium stearate is a slightly flammable white powder with lubricative properties. In pharmaceutical technology it is used in the manufacturing of tablets, to prevent the tablets from sticking to the equipment during the tablet compression process (i.e., when the tablet's substance is pressed into tablet form). • Magnesium sulfite is used in the manufacture of paper (sulfite process). • Magnesium phosphate is used to fireproof wood for construction. • Magnesium hexafluorosilicate is used in mothproofing of textiles.
History The name originates from the Greek word for a district in Thessaly called Magnesia. It is related to magnetite and manganese, which also originated from this area, and required differentiation as separate substances. See manganese for this history. Magnesium is the seventh most abundant element in the Earth's crust by mass and eighth by molarity.[3] It is found in large deposits of magnesite, dolomite, and other minerals, and in mineral waters, where magnesium ion is soluble. In 1618 a farmer at Epsom in England attempted to give his cows water from a well. They refused to drink because of the water's bitter taste. However the farmer noticed that the water seemed to heal scratches and rashes. The fame of Epsom salts spread. Eventually they were recognized to be hydrated magnesium sulfate, MgSO4. The metal itself was first produced in England by Sir Humphry Davy in 1808 using electrolysis of a mixture of magnesia and mercury oxide. Antoine Bussy prepared it in coherent form in 1831. Davy's first suggestion for a name was magnium, but the name magnesium is now used.
6
Magnesium
Sources Ores Although magnesium is found in over 60 minerals, only dolomite, magnesite, brucite, carnallite, talc, and olivine are of commercial importance.
Sea water The Mg2+ cation is the second most abundant cation in sea water (occurring at about 12% of the mass of sodium there), which makes sea water and sea-salt an attractive commercial source of Mg. To extract the magnesium, calcium hydroxide is added to sea water to form magnesium hydroxide precipitate. MgCl2 + Ca(OH)2 → Mg(OH)2 + CaCl2 Magnesium hydroxide is insoluble in water so it can be filtered out, and reacted with hydrochloric acid to obtain concentrated magnesium chloride. Mg(OH)2 + 2 HCl → MgCl2 + 2 H2O From magnesium chloride, electrolysis produces magnesium.
Electrolysis In the United States, magnesium is principally obtained by electrolysis of fused magnesium chloride from brines, wells, and sea water. At the cathode, the Mg2+ ion is reduced by two electrons to magnesium metal: Mg2+ + 2 e− → Mg At the anode, each pair of Cl− ions is oxidised to chlorine gas, releasing two electrons to complete the circuit: 2 Cl− → Cl2 (g) + 2 e− The United States has traditionally been the major world supplier of this metal, supplying 45% of world production even as recently as 1995. Today, the US market share is at 7%, with a single domestic producer left, US Magnesium, a company born from now-defunct Magcorp.[8] As of 2005, China has taken over as the dominant supplier, pegged at 60% world market share, which increased from 4% in 1995. Unlike the above described electrolytic process, China is almost completely reliant on a different method of obtaining the metal from its ores, the silicothermic Pidgeon process (the reduction of the oxide at high temperatures with silicon).
Biological role Due to the important interaction between phosphate and magnesium ions, magnesium ions are essential to the basic nucleic acid chemistry of life, and thus are essential to all cells of all known living organisms. Over 300 enzymes require the presence of magnesium ions for their catalytic action, including all enzymes utilizing or synthesizing ATP, or those which use other nucleotides to synthesize DNA and RNA. ATP exists in cells normally as a chelate of ATP and a magnesium ion. Plants have an additional use for magnesium in that chlorophylls are magnesium-centered porphyrins. Magnesium deficiency in plants causes late-season yellowing between leaf
7
Magnesium veins, especially in older leaves, and can be corrected by applying Epsom salts (which is rapidly leached), or else crushed dolomitic limestone to the soil. Magnesium is a vital component of a healthy human diet. Human magnesium deficiency (including conditions which show few overt symptoms) is relatively common, with only 32% of the United States meeting the RDA-DRI,[9] and has been implicated in the development of a number of human illnesses such as asthma, osteoporosis, and ADHD.[10] Adult human bodies contain about 24 grams of magnesium, with 60% in the skeleton, 39% intracellular (20% in skeletal muscle), and 1% extracellular. Serum levels are typically 0.7 – 1.0 mmol/L. Serum magnesium levels may appear normal even in cases of underlying intracellular deficiency, although no known mechanism maintains a homeostatic level in the blood other than renal excretion of high blood levels. Intracellular magnesium is correlated with intracellular potassium. Magnesium is absorbed in the gastrointestinal tract, Food sources of magnesium with more absorbed when status is lower. In humans, magnesium appears to facilitate calcium absorption. Low and high protein intake inhibit magnesium absorption, and other factors such as phosphate, phytate, and fat affect absorption. Absorbed dietary magnesium is largely excreted through the urine, although most magnesium "administered orally" is excreted through the feces.[11] Magnesium status may be assessed roughly through serum and erythrocyte Mg concentrations and urinary and fecal excretion, but intravenous magnesium loading tests are likely the most accurate and practical in most people.[12] In these tests, magnesium is injected intravenously; a retention of 20% or more indicates deficiency.[13] Other nutrient deficiencies are identified through biomarkers, but none are established for magnesium.[14] Spices, nuts, cereals, coffee, cocoa, tea, and vegetables (especially green leafy ones) are rich sources of magnesium. Observations of reduced dietary magnesium intake in modern Western countries as compared to earlier generations may be related to food refining and modern fertilizers which contain no magnesium.[11] There are a number of magnesium dietary supplements available. Magnesium oxide, one of the most common because it has a high magnesium content per weight, has been reported to be the least bioavailable.[15] [16] Magnesium citrate has been reported as more bioavailable than oxide or amino-acid chelate (glycinate) forms.[17] Excess magnesium in the blood is freely filtered at the kidneys, and for this reason it is difficult to overdose on magnesium from dietary sources alone.[10] With supplements, overdose is possible, however, particularly in people with poor renal function; occasionally, with use of high cathartic doses of magnesium salts, severe hypermagnesemia has been reported to occur even without renal dysfunction.[18] Alcoholism can produce a magnesium deficiency which is easily reversed by oral or parenteral administration, depending on the degree of deficiency.[19]
8
Magnesium
9
Isotopes Magnesium has three stable isotopes: 24Mg, 25Mg, 26Mg. All are present in significant amounts (see table of isotopes above). About 79% of Mg is 24Mg. The isotope 28Mg is radioactive and in the 1950s to 1970s was made commercially by several nuclear power plants for use in scientific experiments. This isotope has a relatively short half-life (21 hours) and so its use was limited by shipping times. 26
Mg has found application in isotopic geology, similar to that of aluminium. 26Mg is a radiogenic daughter product of 26Al, which has a half-life of 717,000 years. Large enrichments of stable 26Mg have been observed in the Ca-Al-rich inclusions of some carbonaceous chondrite meteorites. The anomalous abundance of 26Mg is attributed to the decay of its parent 26Al in the inclusions. Therefore, the meteorite must have formed in the solar nebula before the 26Al had decayed. Hence, these fragments are among the oldest objects in the solar system and have preserved information about its early history. It is conventional to plot 26Mg/24Mg against an Al/Mg ratio. In an isochron dating plot, the Al/Mg ratio plotted is27Al/24Mg. The slope of the isochron has no age significance, but indicates the initial 26Al/27Al ratio in the sample at the time when the systems were separated from a common reservoir.
Precautions Magnesium metal and its alloys are explosive hazards; they are highly flammable in their pure form when molten or in powder or in ribbon form. Burning or molten magnesium metal reacts violently with water. When working with powdered magnesium, safety glasses with welding eye protection are employed, because the bright white light produced by burning magnesium contains ultraviolet light that can permanently damage the retinas of the eyes.[20] Magnesium is capable of reducing water to the highly-flammable hydrogen gas:[21]
The magnesium-bodied Honda RA302 of Jo Schlesser crashes and burns during the 1968 French Grand Prix. Schlesser was killed.
Mg (s) + 2 H2O → Mg(OH)2 (s) + H2 (g) As a result, water cannot be used to extinguish magnesium fires; the hydrogen gas produced will only intensify the fire. Magnesium also reacts with carbon dioxide to form magnesium oxide and carbon: 2 Mg (s) + CO2 → 2 MgO (s) + C (s) Hence, carbon dioxide fire extinguishers cannot be used for extinguishing magnesium fires either.[22] Burning magnesium is usually quenched by using a Class D dry chemical fire extinguisher, or by covering the fire with sand or magnesium foundry flux to remove its air source.
Magnesium
10
External links • WebElements.com – Magnesium [23] • Online Resource for industry professionals
[24]
- Magnesium.com
[25]
• The Magnesium Website – Includes full text papers and textbook chapters by leading magnesium authorities Mildred Seelig, Jean Durlach, Burton M. Altura and Bella T. Altura. Links to over 300 articles discussing magnesium and magnesium deficiency. • Magnesium in Health [26] - Mg12.info
References [1] Bernath, P. F., Black, J. H., & Brault, J. W. (1985). " The spectrum of magnesium hydride (http:/ / bernath. uwaterloo. ca/ media/ 24. pdf)". Astrophysical Journal 298: 375. . [2] Ash, Russell (2005), The Top 10 of Everything 2006: The Ultimate Book of Lists (http:/ / plymouthlibrary. org/ faqelements. htm), Dk Pub, ISBN 0756613213, . [3] (PDF) Abundance and form of the most abundant elements in Earth’s continental crust (http:/ / www. gly. uga. edu/ railsback/ Fundamentals/ ElementalAbundanceTableP. pdf). . Retrieved 2008-02-15. [4] http:/ / www. seafriends. org. nz/ oceano/ seawater. htm#composition. Retrieved Jan. 20, 2009 [5] , http:/ / www. mg12. info [6] DOE Handbook - Primer on Spontaneous Heating and Pyrophoricity (http:/ / www. hss. doe. gov/ nuclearsafety/ ns/ techstds/ standard/ hdbk1081/ hbk1081c. html). U.S. Department of Energy. December 1994. p. 20. DOE-HDBK-1081-94. . [7] Alan A. Luo and Bob R. Powell (2001) (PDF). Tensile and Compressive Creep of Magnesium-Aluminum-Calcium Based Alloys (http:/ / doc. tms. org/ ezmerchant/ prodtms. nsf/ productlookupitemid/ 01-481x-137/ $FILE/ 01-481X-137F. pdf). Materials & Processes Laboratory, General Motors Research & Development Center. . Retrieved 2007-08-21. [8] Vardi, Nathan (February 22 2007). " Man With Many Enemies (http:/ / www. forbes. com/ forbes/ 2002/ 0722/ 044_print. html)". Forbes.com. . Retrieved 2006-06-26. [9] " Lack Energy? Maybe It's Your Magnesium Level (http:/ / www. ars. usda. gov/ is/ AR/ archive/ may04/ energy0504. htm?pf=1)". United States Department of Agriculture. . Retrieved 2008-09-18. Last paragraph [10] University of Maryland Medical Center. Magnesium (http:/ / www. umm. edu/ altmed/ articles/ magnesium-000313. htm) [11] Wester PO (May 1987). " Magnesium (http:/ / www. ajcn. org/ cgi/ pmidlookup?view=long& pmid=3578120)". Am. J. Clin. Nutr. 45 (5 Suppl): 1305–12. PMID 3578120. . [12] Arnaud MJ (June 2008). "Update on the assessment of magnesium status". Br. J. Nutr. 99 Suppl 3: S24–36. doi: 10.1017/S000711450800682X (http:/ / dx. doi. org/ 10. 1017/ S000711450800682X). PMID 18598586. [13] Rob PM, Dick K, Bley N, et al. (October 1999). " Can one really measure magnesium deficiency using the short-term magnesium loading test? (http:/ / www. blackwell-synergy. com/ openurl?genre=article& sid=nlm:pubmed& issn=0954-6820& date=1999& volume=246& issue=4& spage=373)". J. Intern. Med. 246 (4): 373–8. doi: 10.1046/j.1365-2796.1999.00580.x (http:/ / dx. doi. org/ 10. 1046/ j. 1365-2796. 1999. 00580. x). PMID 10583708. . [14] Franz KB (December 2004). " A functional biological marker is needed for diagnosing magnesium deficiency (http:/ / www. jacn. org/ cgi/ pmidlookup?view=long& pmid=15637224)". J Am Coll Nutr 23 (6): 738S–41S. PMID 15637224. . [15] Firoz M, Graber M (December 2001). "Bioavailability of US commercial magnesium preparations". Magnes Res 14 (4): 257–62. PMID 11794633. [16] Lindberg JS, Zobitz MM, Poindexter JR, Pak CY (February 1990). "Magnesium bioavailability from magnesium citrate and magnesium oxide". J Am Coll Nutr 9 (1): 48–55. PMID 2407766. [17] Walker AF, Marakis G, Christie S, Byng M (September 2003). " Mg citrate found more bioavailable than other Mg preparations in a randomised, double-blind study (http:/ / www. john-libbey-eurotext. fr/ medline. md?issn=0953-1424& vol=16& iss=3& page=183)". Magnes Res 16 (3): 183–91. PMID 14596323. . [18] Kontani M, Hara A, Ohta S, Ikeda T (2005). "Hypermagnesemia induced by massive cathartic ingestion in an elderly woman without pre-existing renal dysfunction". Intern. Med. 44 (5): 448–52. doi: 10.2169/internalmedicine.44.448 (http:/ / dx. doi. org/ 10. 2169/ internalmedicine. 44. 448). PMID 15942092. [19] AJ Giannini. Drugs of Abuse--Second Edition. Los Angeles, Physicians Management Information Co., 1997. [20] " Science Safety: Chapter 8 (http:/ / www. edu. gov. mb. ca/ k12/ docs/ support/ scisafe/ chapter8. html)". Government of Manitoba. . Retrieved 2007-08-21.
Magnesium [21] " Chemistry : Periodic Table : magnesium : chemical reaction data (http:/ / www. webelements. com/ webelements/ elements/ text/ Mg/ chem. html)". webelements.com. . Retrieved 2006-06-26. [22] " Demo Lab: Reaction Of Magnesium Metal With Carbon Dioxide (http:/ / www. ilpi. com/ genchem/ demo/ co2mg/ )". . Retrieved 2006-06-26. [23] [24] [25] [26]
http:/ / www. webelements. com/ magnesium/ http:/ / www. magnesium. com http:/ / www. mgwater. com/ index. shtml http:/ / www. mg12. info
11
Article Sources and Contributors
Article Sources and Contributors Magnesium Source: http://en.wikipedia.org/w/index.php?oldid=307135069 Contributors: 130.94.122.xxx, 17Drew, 19.168, 4-409r-0, 65.68.87.xxx, 7, 8472, A. di M., ADM003, Aaeamdar, Abeg92, Abrutt, Ace11423, Adashiel, Addshore, Aditya, Aff123a, Ahoerstemeier, Aitias, Ajaxkroon, Aksi great, Alansohn, Ale jrb, Alexburke, Alexf, AlexiusHoratius, Alphachimp, Amalthea, Amwyll Rwden, Andre Engels, Andres, Andrewpmk, AndyVolykhov, Antandrus, Apparition11, Arakunem, Archimerged, Arcyqwerty, Arjun01, Arundhati bakshi, Ascidian, Astrowob, Atchius, Atlant, Austoria43, Avi saig, AzaToth, Badgernet, Bbatsell, BeautifulMachine, Beetstra, Bennybp, Bently34, Bhadani, Bhound89, Bigtop, BillFlis, Biochemnick, Bjweeks, BlueEarth, Bluebec, Bobblewik, Bobo192, Bogey97, Bold Clone, Bomac, BoomerAB, Bork, Brian0918, Brinerustle, Bryan Derksen, Bsimmons666, Bubbha, CYD, Cadmium, Caltas, Can't sleep, clown will eat me, CanadianLinuxUser, Cardil, CardinalDan, Carinemily, Carnildo, Catbar, Caulde, Ceyockey, Cflm001, Chadlupkes, Charles Gaudette, Chickyfuzz123, Chris 73, Chrislk02, Chriswiki, Chromaticity, Church of emacs, Cireshoe, Citicat, ClanCC, Cometstyles, Conversion script, Coolbeans101, Coolio213, Coppertwig, Corpx, Costelld, Craftyminion, Cremepuff222, Cssiitcic, Cureden, DVD R W, Dacar92, Dale Arnett, Damicatz, Dan100, Darrien, Davewho2, David Latapie, Db099221, Deglr6328, Dekisugi, Delirium, Delta G, Dillin268, Dirkbb, Discospinster, Dmoskva, DonSiano, DoubleBlue, Dr. Morbius, DrBob, Dreadstar, DuO, Duk, Dwmyers, Dysepsion, EL Willy, Eaolson, Early account, Eastlaw, EddEdmondson, Edgar181, Edsanville, El C, Element16, Ellsworth, Emc2, Enok Walker, Enviroboy, Enzofroilan, Epbr123, Epo, Eric-Wester, Eric119, Ericd, Everyking, Evil saltine, Excirial, Faeflora, Femto, Fenrir, Fieldday-sunday, Figureskatingfan, Finalbastion, Flosseveryday, FocalPoint, Fonzy, Frankenpuppy, FreplySpang, Fuzbaby, GT5162, Gajakk, Galvotec, Gamera2, Gargaj, Gaurav1146, Gene Nygaard, Geniac, Ggonnell, Giftlite, Gigs, Gilliam, Gman124, Gorbb, GrahamColm, Grant M, Grantus4504, Grendelkhan, Gunnar Hendrich, Gwernol, Gymmery, Gzkn, Gökhan, Hadal, Hak-kâ-ngìn, HalJor, Hantzen, Harlequence, Hellbus, HenryLi, Herbee, HereToHelp, Heron, HexaChord, HonztheBusDriver, Hotcrocodile, Hottstuff111, Huntthetroll, Hut 8.5, II MusLiM HyBRiD II, IRP, Ian Spackman, Icairns, Ich, Igoldste, Ikiroid, Im not maaad, ImperfectlyInformed, Ioeth, Iridescence, Iridescent, Irishguy, Island, IsmAvatar, Ixfd64, J.delanoy, JForget, Jaraalbe, Jaredroberts, Jasz, Jaybo007, Jeff G., Jennavecia, Jeronimo, JesseW, Jessepmullan, Jeversol, Jj137, John, Johner, Jonathan Hall, Jose77, Joyous!, Jqt, Justforasecond, Kaare, Karl-Henner, Karlhahn, Karn, Keenanpepper, Keilana, Kelly Martin, Kf4bdy, Kilo-Lima, King Lopez, Kingpin13, KnowledgeOfSelf, Kpjas, Kris Schnee, Kuru, Kurykh, Kwamikagami, LeaveSleaves, Lec CRP1, Lee J Haywood, Leila3, Lifung, LittleOldMe, Loren.wilton, LorenzoB, Lradrama, LuigiManiac, Luigifan07, Luna Santin, Luxdormiens, MER-C, MZMcBride, Ma8thew, Makemi, Malcolm Farmer, Mani1, Marc Kupper, Marlith, Martinman11, Materialscientist, Mattcain, Mav, Maximus Rex, May0208, Mdf, Melchoir, Mentifisto, Merope, Mgimpel, Michaelas10, Mike Gale, Mikiemike, Minesweeper, Miss Madeline, Mixwell, Mmm, Modulatum, Moe Epsilon, Monkeyman, Moonasha, Mormor1, Mr0t1633, Mxn, NCurse, NEIL4737, NHRHS2010, NSK Nikolaos S. 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Image Sources, Licenses and Contributors file:hexagonal.svg Source: http://en.wikipedia.org/w/index.php?title=File:Hexagonal.svg License: BSD Contributors: Original uploader was Danieljamesscott at en.wikipedia file:Electron shell 012 Magnesium.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_012_Magnesium.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 file:Magnesium crystals.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Magnesium_crystals.jpg License: unknown Contributors: User:Warut Image:Bundesarchiv Bild 102-12062, Wasserreiter mit Magnesiumfackeln.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Bundesarchiv_Bild_102-12062,_Wasserreiter_mit_Magnesiumfackeln.jpg License: unknown Contributors: Gamsbart, Mattes, NSK Nikolaos S. Karastathis, Raven1977 Image:Magnesium-products.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Magnesium-products.jpg License: unknown Contributors: Firetwister, Warut, 3 anonymous edits Image:Magnesium Sparks.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Magnesium_Sparks.jpg License: GNU Free Documentation License Contributors: Hiroaki Nakamura Image:FoodSourcesOfMagnesium.jpg Source: http://en.wikipedia.org/w/index.php?title=File:FoodSourcesOfMagnesium.jpg License: Public Domain Contributors: Peggy Greb Image:Schlesser.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Schlesser.jpg License: unknown Contributors: Lec CRP1
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12
Aluminium
1
Aluminium magnesium ← aluminium → siliconB ↑ Al ↓ Ga
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
13Al Periodic table
Appearance grey
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
1482
1632
1817
2054
2364
2790
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1816.7 kJ·mol−1 3rd: 2744.8 kJ·mol−1Atomic radiusCovalent radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal
Aluminium
2
conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusPoisson ratioMohs hardnessVickers hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of aluminium iso 26
Al
27
Al
N.A. syn
100%
half-life 7.17×105y
DM
DE (MeV)
DP
β+
1.17
26
ε
-
26
γ
1.8086
-
Mg Mg
27
Al is stable with 14 neutron
aluminium, Al, 13 poor metal13, 3, p26.9815386(13) g·mol−1 [Ne] 3s2 3p1 2, 8, 3 (Image) solid 2.70 g·cm−3 2.375 g·cm−3 933.47 K,660.32 °C,1220.58 °F 2792 K,2519 °C,4566 °F 10.71 kJ·mol−1 294.0 kJ·mol−1 (25 °C) 24.200 J·mol−1·K−13, 2[1] , 1[2] (amphoteric oxide) 1.61 (Pauling scale) 1st: 577.5 kJ·mol−1143 pm121±4 pm face-centered cubic paramagnetic[3] (20 °C) 28.2 nΩ·m (300 K) 237 W·m−1·K−1 (25 °C) 23.1 µm·m−1·K−1 (r.t.) (rolled) 5,000 m·s−1 70 GPa 26 GPa 76 GPa 0.35 2.75 167 MPa 245 MPa 7429-90-5 Aluminium ( ˌæljʊˈmɪniəm Wikipedia:Media helpFile:En-uk-aluminium1.ogg) or aluminum ( /əˈluːmɪnəm/ Wikipedia:Media helpFile:En-uk-aluminum.ogg, see spelling below) is a silvery white and ductile member of the boron group of chemical elements. It has the symbol Al; its atomic number is 13. It is not soluble in water under normal circumstances. Aluminium is the most abundant metal in the Earth's crust, and the third most abundant element therein, after oxygen and silicon. It makes up about 8% by weight of the Earth’s solid surface. Aluminium is too reactive chemically to occur in nature as a free metal. Instead, it is found combined in over 270 different minerals.[4] The chief source of aluminium is bauxite ore. Aluminium is remarkable for its ability to resist corrosion due to the phenomenon of passivation and for the metal's low density. Structural components made from aluminium and its alloys are vital to the aerospace industry and very important in other areas of transportation and building. Its reactive nature makes it useful as a catalyst or additive in chemical mixtures, including being used in ammonium nitrate explosives to enhance blast power.
Characteristics Aluminium is a soft, durable, lightweight, malleable metal with appearance ranging from silvery to dull grey, depending on the surface roughness. Aluminium is nonmagnetic and nonsparking. It is also insoluble in alcohol, though it can be soluble in water in certain forms. The yield strength of pure aluminium is 7–11 MPa, while aluminium alloys have yield strengths ranging from 200 MPa to 600 MPa.[5] Aluminium has about one-third the density and stiffness of steel. It is ductile, and easily machined, cast, and extruded. Corrosion resistance can be excellent due to a thin surface layer of aluminium oxide that forms when the metal is exposed to air, effectively preventing further oxidation. The strongest aluminium alloys are less corrosion resistant due to galvanic reactions with alloyed copper.[5] This corrosion resistance is also often greatly reduced when many aqueous salts are present however, particularly in the presence of dissimilar metals.
Aluminium Aluminium atoms are arranged in a face-centered cubic (fcc) structure. Aluminium has a stacking-fault energy of approximately 200 mJ/m².[6] Aluminium is one of the few metals that retain full silvery reflectance in finely powdered form, making it an important component of silver paints. Aluminium mirror finish has the highest reflectance of any metal in the 200–400 nm (UV) and the 3000–10000 nm (far IR) regions, while in the 400–700 nm visible range it is slightly outdone by tin and silver and in the 700–3000 (near IR) by silver, gold, and copper.[7] Aluminium is a good thermal and electrical conductor, by weight better than copper. Aluminium is capable of being a superconductor, with a superconducting critical temperature of 1.2 kelvin and a critical magnetic field of about 100 gauss.[8]
Isotopes Aluminium has nine isotopes, whose mass numbers range from 23 to 30. Only 27Al (stable isotope) and 26Al (radioactive isotope, t1/2 = 7.2 × 105 y) occur naturally; however, 27Al has a natural abundance of 99.9+ %. 26Al is produced from argon in the atmosphere by spallation caused by cosmic-ray protons. Aluminium isotopes have found practical application in dating marine sediments, manganese nodules, glacial ice, quartz in rock exposures, and meteorites. The ratio of 26Al to 10Be has been used to study the role of transport, deposition, sediment storage, burial times, and erosion on 105 to 106 year time scales.[9] Cosmogenic 26Al was first applied in studies of the Moon and meteorites. Meteoroid fragments, after departure from their parent bodies, are exposed to intense cosmic-ray bombardment during their travel through space, causing substantial 26Al production. After falling to Earth, atmospheric shielding protects the meteorite fragments from further 26Al production, and its decay can then be used to determine the meteorite's terrestrial age. Meteorite research has also shown that 26Al was relatively abundant at the time of formation of our planetary system. Most meteoriticists believe that the energy released by the decay of 26Al was responsible for the melting and differentiation of some asteroids after their formation 4.55 billion years ago.[10]
Natural occurrence In the Earth's crust, aluminium is the most abundant (8.3% by weight) metallic element and the third most abundant of all elements (after oxygen and silicon).[11] Because of its strong affinity to oxygen, however, it is almost never found in the elemental state; instead it is found in oxides or silicates. Feldspars, the most common group of minerals in the Earth's crust, are aluminosilicates. Native aluminium metal can be found as a minor phase in low oxygen fugacity environments, such as the interiors of certain volcanoes.[12] It also occurs in the minerals beryl, cryolite, garnet, spinel and turquoise.[11] Impurities in Al2O3, such as chromium or cobalt yield the gemstones ruby and sapphire, respectively. Pure Al2O3, known as Corundum, is one of the hardest materials known.[11] Although aluminium is an extremely common and widespread element, the common aluminium minerals are not economic sources of the metal. Almost all metallic aluminium is produced from the ore bauxite (AlOx(OH)3-2x). Bauxite occurs as a weathering product of low iron and silica bedrock in tropical climatic conditions.[13] Large deposits of bauxite occur in Australia, Brazil, Guinea and Jamaica but the primary mining areas for the ore are in Ghana, Indonesia, Jamaica, Russia and Surinam.[14] Smelting of the ore mainly occurs in Australia, Brazil, Canada, Norway, Russia and the United States. Because smelting is an
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Aluminium energy-intensive process, regions with excess natural gas supplies (such as the United Arab Emirates) are becoming aluminium refiners.
Production and refinement Although aluminium is the most abundant metallic element in the Earth's crust (believed to be 7.5 to 8.1 percent), it is rare in its free form, occurring in oxygen-deficient environments such as volcanic mud, and it was once considered a precious metal more valuable than gold. Napoleon III, emperor of France, is reputed to have given a banquet where the most honoured guests were given aluminium utensils, while the other guests had to make do with gold.[15] [16] The Washington Monument was completed, with the 100 ounce (2.8 kg) aluminium capstone being put in place on December 6, 1884, in an elaborate dedication ceremony. It was the largest single piece of aluminium cast at the time. At that time, aluminium was as expensive as silver.[17] Aluminium has been produced in commercial quantities for just over 100 years. Aluminium is a strongly reactive metal that forms a high-energy chemical bond with oxygen. Compared to most other metals, it is difficult to extract from ore, such as bauxite, due to the energy required to reduce aluminium oxide (Al2O3). For example, direct reduction with carbon, as is used to produce iron, is not chemically possible, since aluminium is a stronger reducing agent than carbon. Aluminium oxide has a melting point of about 2,000 °C. Therefore, it must be Bauxite extracted by electrolysis. In this process, the aluminium oxide is dissolved in molten cryolite and then reduced to the pure metal. The operational temperature of the reduction cells is around 950 to 980 °C. Cryolite is found as a mineral in Greenland, but in industrial use it has been replaced by a synthetic substance. Cryolite is a chemical compound of aluminium, sodium, and calcium fluorides: (Na3AlF6). The aluminium oxide (a white powder) is obtained by refining bauxite in the Bayer process of Karl Bayer. (Previously, the Deville process was the predominant refining technology.) The electrolytic process replaced the Wöhler process, which involved the reduction of anhydrous aluminium chloride with potassium. Both of the electrodes used in the electrolysis of aluminium oxide are carbon. Once the refined alumina is dissolved in the electrolyte, its ions are free to move around. The reaction at the cathode (negative electrode) is Al3+ + 3 e− → Al Here the aluminium ion is being reduced (electrons are added). The aluminium metal then sinks to the bottom and is tapped off, usually cast into large blocks called aluminium billets for further processing. At the anode (positive electrode), oxygen is formed: 2 O2− → O2 + 4 e− This carbon anode is then oxidized by the oxygen, releasing carbon dioxide. O2 + C → CO2
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Aluminium The anodes in a reduction cell must therefore be replaced regularly, since they are consumed in the process. Unlike the anodes, the cathodes are not oxidized because there is no oxygen present, as the carbon cathodes are protected by the liquid aluminium inside the cells. Nevertheless, cathodes do erode, mainly due to electrochemical processes and metal movement. After five to ten years, depending on the current used in the electrolysis, a cell has to be rebuilt because of cathode wear. Aluminium electrolysis with the Hall-Héroult process consumes a lot of energy, but alternative processes were always found to be less viable economically and/or ecologically. The worldwide average specific energy consumption is approximately 15±0.5 kilowatt-hours per kilogram of aluminium produced (52 to 56 MJ/kg). The most modern smelters achieve approximately 12.8 kW·h/kg (46.1 MJ/kg). (Compare this to the heat of World production trend of aluminium reaction, 31 MJ/kg, and the Gibbs free energy of reaction, 29 MJ/kg.) Reduction line currents for older technologies are typically 100 to 200 kA; state-of-the-art smelters[18] operate at about 350 kA. Trials have been reported with 500 kA cells. Electric power represents about 20% to 40% of the cost of producing aluminium, depending on the location of the smelter. Smelters tend to be situated where electric power is both plentiful and inexpensive, such as South Africa, Ghana, the South Island of New Zealand, Australia, the People's Republic of China, the Middle East, Russia, Quebec and British Columbia in Canada, and Iceland.[19] In 2005, the People's Republic of China was the top producer of aluminium with almost a one-fifth world share, followed by Russia, Canada, and the USA, reports the British Geological Survey. Over the last 50 years, Australia has become a major producer of bauxite ore and a major producer and Aluminium output in 2005 [20] exporter of alumina. Australia produced 62 million tonnes of bauxite in 2005. The Australian deposits have some refining problems, some being high in silica but have the advantage of being shallow and relatively easy to mine.[21]
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Aluminium
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Recycling Aluminium is 100% recyclable without any loss of its natural qualities. Recovery of the metal via recycling has become an important facet of the aluminium industry. Recycling involves melting the scrap, a process that requires only five percent of the energy used to produce aluminium from ore. However, a significant part (up to 15% of the input material) is lost as dross (ash-like oxide).[22] Recycling was a low-profile activity until the late 1960s, when the growing use of aluminium beverage cans brought it to the public awareness.
Aluminium Recycling Code
In Europe aluminium experiences high rates of recycling, ranging from 42% of beverage cans, 85% of construction materials and 95% of transport vehicles.[23] Recycled aluminium is known as secondary aluminium, but maintains the same physical properties as primary aluminium. Secondary aluminium is produced in a wide range of formats and is employed in 80% of the alloy injections. Another important use is for extrusion. White dross from primary aluminium production and from secondary recycling operations still contains useful quantities of aluminium which can be extracted industrially.[24] The process produces aluminium billets, together with a highly complex waste material. This waste is difficult to manage. It reacts with water, releasing a mixture of gases (including, among others, hydrogen, acetylene, and ammonia) which spontaneously ignites on contact with air;[25] contact with damp air results in the release of copious quantities of ammonia gas. Despite these difficulties, however, the waste has found use as a filler in asphalt and concrete.[26]
Chemistry Oxidation state one AlH is produced when aluminium is heated in an atmosphere of hydrogen. Al2O is made by heating the normal oxide, Al2O3, with silicon at 1800 °C in a vacuum.[27] Al2S can be made by heating Al2S3 with aluminium shavings at 1300 °C in a vacuum.[27] It quickly disproportionates to the starting materials. The selenide is made in a parallel manner. AlF, AlCl and AlBr exist in the gaseous phase when the tri-halide is heated with aluminium. Aluminium halides usually exist in the form AlX3. e.g. AlF3, AlCl3, AlBr3, AlI3 etc.[27]
Aluminium
Oxidation state two Aluminium monoxide, AlO, has been detected in the gas phase after explosion[28] and in stellar absorption spectra.[29]
Oxidation state three Fajans' rules show that the simple trivalent cation Al3+ is not expected to be found in anhydrous salts or binary compounds such as Al2O3. The hydroxide is a weak base and aluminium salts of weak acids, such as carbonate, can't be prepared. The salts of strong acids, such as nitrate, are stable and soluble in water, forming hydrates with at least six molecules of water of crystallization. Aluminium hydride, (AlH3)n, can be produced from trimethylaluminium and an excess of hydrogen. It burns explosively in air. It can also be prepared by the action of aluminium chloride on lithium hydride in ether solution, but cannot be isolated free from the solvent. Alumino-hydrides of the most electropositive elements are known, the most useful being lithium aluminium hydride, Li[AlH4]. It decomposes into lithium hydride, aluminium and hydrogen when heated, and is hydrolysed by water. It has many uses in organic chemistry, particularly as a reducing agent. The aluminohalides have a similar structure. Aluminium hydroxide may be prepared as a gelatinous precipitate by adding ammonia to an aqueous solution of an aluminium salt. It is amphoteric, being both a very weak acid, and forming aluminates with alkalis. It exists in various crystalline forms. Aluminium carbide, Al4C3 is made by heating a mixture of the elements above 1000 °C. The pale yellow crystals have a complex lattice structure, and react with water or dilute acids to give methane. The acetylide, Al2(C2)3, is made by passing acetylene over heated aluminium. Aluminium nitride, AlN, can be made from the elements at 800 °C. It is hydrolysed by water to form ammonia and aluminium hydroxide. Aluminium phosphide, AlP, is made similarly, and hydrolyses to give phosphine. Aluminium oxide, Al2O3, occurs naturally as corundum, and can be made by burning aluminium in oxygen or by heating the hydroxide, nitrate or sulfate. As a gemstone, its hardness is only exceeded by diamond, boron nitride, and carborundum. It is almost insoluble in water. Aluminium sulfide, Al2S3, may be prepared by passing hydrogen sulfide over aluminium powder. It is polymorphic. Aluminium iodide, AlI3, is a dimer with applications in organic synthesis. Aluminium fluoride, AlF3, is made by treating the hydroxide with HF, or can be made from the elements. It consists of a giant molecule which sublimes without melting at 1291 °C. It is very inert. The other trihalides are dimeric, having a bridge-like structure. Aluminium fluoride/water complexes: When aluminium and fluoride are together in aqueous solution, they readily form complex ions such as AlF(H2O)5+2, AlF3(H2O)30, AlF6−3. Of these, AlF6−3 is the most stable. This is explained by the fact that aluminium and fluoride, which are both very compact ions, fit together just right to form the octahedral aluminium hexafluoride complex. When aluminium and fluoride are together in water in a 1:6 molar ratio, AlF6−3 is the most common form, even in rather low concentrations. Organo-metallic compounds of empirical formula AlR3 exist and, if not also giant molecules, are at least dimers or trimers. They have some uses in organic synthesis, for instance trimethylaluminium. Analysis
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Aluminium
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The presence of aluminium can be detected in qualitative analysis using aluminon.
Applications General use Aluminium is the most widely used non-ferrous metal.[30] Global production of aluminium in 2005 was 31.9 million tonnes. It exceeded that of any other metal except iron (837.5 million tonnes).[31] Relatively pure aluminium is encountered only when corrosion resistance and/or workability is more important than strength or hardness. A thin layer of aluminium can be deposited onto a flat surface by physical vapor deposition or (very infrequently) chemical vapor deposition or other chemical means to form optical coatings and mirrors. When so deposited, a fresh, pure aluminium film serves as a good reflector (approximately 92%) of visible light and an excellent reflector (as much as 98%) of medium and far infrared radiation. Pure aluminium has a low tensile strength, but when combined with thermo-mechanical processing, aluminium alloys display a marked improvement in mechanical properties, especially when tempered. Aluminium alloys form vital components of aircraft and rockets as a result of their high strength-to-weight ratio. Aluminium readily forms alloys with many elements such as copper, zinc, magnesium, manganese and silicon (e.g., duralumin). Today, almost all bulk metal materials that are referred to loosely as "aluminium," are actually alloys. For example, the common aluminium foils are alloys of 92% to 99% aluminium.[32] Some of the many uses for aluminium metal are in: • Transportation (automobiles, aircraft, trucks, railway cars, marine vessels, bicycles etc.) as sheet, tube, castings etc. • Packaging (cans, foil, etc.) • Construction (windows, doors, siding, building wire, etc.) • A wide range of household items, from cooking utensils to baseball bats, watches[33] and notebook computers (Apple) • Street lighting poles, sailing ship masts, walking poles etc.
Household aluminium foil
• Outer shells of consumer electronics, also cases for equipment e.g. photographic equipment. • Electrical transmission lines for power distribution • MKM steel and Alnico magnets • Super purity aluminium (SPA, 99.980% to 99.999% Al), used in electronics and CDs. • Heat sinks for electronic appliances such as transistors and CPUs. • Substrate material of metal-core copper clad laminates used in high brightness LED lighting.
Aluminium-bodied Austin "A40 Sports"(circa 1951)
• Powdered aluminium is used in paint, and in pyrotechnics such as solid rocket fuels and thermite.
Aluminium
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Aluminium compounds • Aluminium ammonium sulfate ([Al(NH4)](SO4)2), ammonium alum is used as a mordant, in water purification and sewage treatment, in paper production, as a food additive, and in leather tanning. • Aluminium acetate is a salt used in solution as an astringent. • Aluminium borate (Al2O3 B2O3) is used in the production of glass and ceramic.
Aluminium slab being transported from the smelters
• Aluminium borohydride (Al(BH4)3) is used as an additive to jet fuel. • Aluminium bronze (CuAl5) • Aluminium chloride (AlCl3) is used: in paint manufacturing, in antiperspirants, in petroleum refining and in the production of synthetic rubber. • Aluminium chlorohydrate is used as an antiperspirant and in the treatment of hyperhidrosis. • Aluminium fluorosilicate (Al2(SiF6)3) is used in the production of synthetic gemstones, glass and ceramic. • Aluminium hydroxide (Al(OH)3) is used: as an antacid, as a mordant, in water purification, in the manufacture of glass and ceramic and in the waterproofing of fabrics. • Aluminium oxide (Al2O3), alumina, is found naturally as corundum (rubies and sapphires), emery, and is used in glass making. Synthetic ruby and sapphire are used in lasers for the production of coherent light. Used as a refractory, essential for the production of high pressure sodium lamps. • Aluminium phosphate (AlPO4) is used in the manufacture: of glass and ceramic, pulp and paper products, cosmetics, paints and varnishes and in making dental cement. • Aluminium sulfate (Al2(SO4)3) is used: in the manufacture of paper, as a mordant, in a fire extinguisher, in water purification and sewage treatment, as a food additive, in fireproofing, and in leather tanning. • Aqueous Aluminium ions (such as found in aqueous Aluminium Sulfate) are use to treat against fish parasites such as Gyrodactylus salaris. • In many vaccines, certain aluminium salts serve as an immune adjuvant (immune response booster) to allow the protein in the vaccine to achieve sufficient potency as an immune stimulant.
Aluminium
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Aluminium alloys in structural applications Aluminium alloys with a wide range of properties are used in engineering structures. Alloy systems are classified by a number system (ANSI) or by names indicating their main alloying constituents (DIN and ISO). The strength and durability of aluminium alloys vary widely, not only as a result of the components of the specific alloy, but also as a result of heat treatments and manufacturing processes. A lack of knowledge of these aspects has from time to time led to improperly designed structures and gained aluminium a bad reputation. (See main article) One important structural limitation of aluminium alloys is their fatigue strength. Unlike steels, aluminium alloys have no well-defined fatigue limit, meaning that fatigue failure will eventually occur under even very small cyclic loadings. This implies that engineers must assess these loads and design for a fixed life rather than an infinite life.
Aluminium foam
Another important property of aluminium alloys is their sensitivity to heat. Workshop procedures involving heating are complicated by the fact that aluminium, unlike steel, will melt without first glowing red. Forming operations where a blow torch is used therefore requires some expertise, since no visual signs reveal how close the material is to melting. Aluminium alloys, like all structural alloys, also are subject to internal stresses following heating operations such as welding and casting. The problem with aluminium alloys in this regard is their low melting point, which make them more susceptible to distortions from thermally induced stress relief. Controlled stress relief can be done during manufacturing by heat-treating the parts in an oven, followed by gradual cooling—in effect annealing the stresses. The low melting point of aluminium alloys has not precluded their use in rocketry; even for use in constructing combustion chambers where gases can reach 3500 K. The Agena upper stage engine used a regeneratively cooled aluminium design for some parts of the nozzle, including the thermally critical throat region.
Household wiring Compared to copper, aluminium has about 65% of the electrical conductivity by volume, although 200% by weight. Traditionally copper is used as household wiring material. In the 1960s aluminium was considerably cheaper than copper, and so was introduced for household electrical wiring in the United States, even though many fixtures had not been designed to accept aluminium wire. In some cases the greater coefficient of thermal expansion of aluminium causes the wire to expand and contract relative to the dissimilar metal screw connection, eventually loosening the connection. Also, pure aluminium has a tendency to creep under steady sustained pressure (to a greater degree as the temperature rises), again loosening the connection. Finally, Galvanic corrosion from the dissimilar metals increased the electrical resistance of the connection.
Aluminium All of this resulted in overheated and loose connections, and this in turn resulted in fires. Builders then became wary of using the wire, and many jurisdictions outlawed its use in very small sizes in new construction. Eventually, newer fixtures were introduced with connections designed to avoid loosening and overheating. The first generation fixtures were marked "Al/Cu" and were ultimately found suitable only for copper-clad aluminium wire, but the second generation fixtures, which bear a "CO/ALR" coding, are rated for unclad aluminium wire. To adapt older assemblies, workers forestall the heating problem using a properly-done crimp of the aluminium wire to a short "pigtail" of copper wire. Today, new alloys, designs, and methods are used for aluminium wiring in combination with aluminium termination.
History Ancient Greeks and Romans used aluminium salts as dyeing mordants and as astringents for dressing wounds; alum is still used as a styptic. In 1761 Guyton de Morveau suggested calling the base alum alumine. In 1808, Humphry Davy identified the existence of a metal base of alum, which he at first termed alumium and later aluminum (see Etymology section, below). The metal was first produced in 1825 (in an impure form) by Danish physicist and chemist Hans Christian Ørsted. He reacted anhydrous aluminium chloride with potassium amalgam and yielded a lump of metal looking similar to tin.[34] Friedrich Wöhler was aware of these experiments and cited them, but after redoing the experiments of Ørsted he concluded that this metal was pure potassium. He conducted a similar experiment in 1827 by mixing anhydrous aluminium chloride with potassium and yielded aluminium.[34] Wöhler is The statue of the Anteros (commonly mistaken for either The Angel of generally credited with isolating aluminium (Latin Christian Charity or Eros) in Piccadilly alumen, alum), but also Ørsted can be listed as its Circus London, was made in 1893 and discoverer.[35] Further, Pierre Berthier discovered is one of the first statues to be cast in aluminium in bauxite ore and successfully extracted aluminium. it.[36] Frenchman Henri Etienne Sainte-Claire Deville improved Wöhler's method in 1846, and described his improvements in a book in 1859, chief among these being the substitution of sodium for the considerably more expensive potassium. (Note: The title of Deville's book is De l'aluminium, ses propriétés, sa fabrication [37] (Paris, 1859). Deville likely also conceived the idea of the electrolysis of aluminium oxide dissolved in cryolite; however, Charles Martin Hall and Paul Héroult might have developed the more practical process after Deville.) Before the Hall-Héroult process was developed, aluminium was exceedingly difficult to extract from its various ores. This made pure aluminium more valuable than gold. Bars of aluminium were exhibited alongside the French crown jewels at the Exposition Universelle of 1855, and Napoleon III was said to have reserved a set of aluminium dinner plates for his
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Aluminium most honoured guests. Aluminium was selected as the material to be used for the apex of the Washington Monument in 1884, a time when one ounce (30 grams) cost the daily wage of a common worker on the project;[38] aluminium was about the same value as silver. The Cowles companies supplied aluminium alloy in quantity in the United States and England using smelters like the furnace of Carl Wilhelm Siemens by 1886.[39] Charles Martin Hall of Ohio in the U.S. and Paul Héroult of France independently developed the Hall-Héroult electrolytic process that made extracting aluminium from minerals cheaper and is now the principal method used worldwide. The Hall-Heroult process cannot produce Super Purity Aluminium directly. Hall's process,[40] in 1888 with the financial backing of Alfred E. Hunt, started the Pittsburgh Reduction Company today known as Alcoa. Héroult's process was in production by 1889 in Switzerland at Aluminium Industrie, now Alcan, and [41] at British Aluminium, now Luxfer Group and Alcoa, by 1896 in Scotland. By 1895 the metal was being used as a building material as far away as Sydney, Australia in the dome of the Chief Secretary's Building. Many navies use an aluminium superstructure for their vessels, however, the 1975 fire aboard USS Belknap that gutted her aluminium superstructure, as well as observation of battle damage to British ships during the Falklands War, led to many navies switching to all steel superstructures. The Arleigh Burke class was the first such U.S. ship, being constructed entirely of steel. In 2008 the price of aluminium peaked at $1.45/lb in July but dropped to $0.7/lb by December.[42]
Etymology Nomenclature history The earliest citation given in the Oxford English Dictionary for any word used as a name for this element is alumium, which British chemist and inventor Humphry Davy employed in 1808 for the metal he was trying to isolate electrolytically from the mineral alumina. The citation is from his journal Philosophical Transactions: "Had I been so fortunate as..to have procured the metallic substances I was in search of, I should have proposed for them the names of silicium, alumium, zirconium, and glucium."[43] By 1812, Davy had settled on aluminum. He wrote in the journal Chemical Philosophy: "As yet Aluminum has not been obtained in a perfectly free state."[44] But the same year, an anonymous contributor to the Quarterly Review, a British political-literary journal, objected to aluminum and proposed the name aluminium, "for so we shall take the liberty of writing the word, in preference to aluminum, which has a less classical sound."[45] The -ium suffix had the advantage of conforming to the precedent set in other newly discovered elements of the time: potassium, sodium, magnesium, calcium, and strontium (all of which Davy had isolated himself). Nevertheless, -um spellings for elements were not unknown at the time, as for example platinum, known to Europeans since the sixteenth century, molybdenum, discovered in 1778, and tantalum, discovered in 1802. The -um suffix on the other hand, has the advantage of being more consistent with the universal spelling alumina for the oxide, as lanthana is the oxide of lanthanum, and magnesia, ceria, and thoria are the oxides of magnesium, cerium, and thorium respectively.
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Aluminium The spelling used throughout the 19th century by most U.S. chemists ended in -ium, but common usage is less clear.[46] The -um spelling is used in the Webster's Dictionary of 1828, as it was in 1892 when Charles Martin Hall published an advertising handbill for his new electrolytic method of producing the metal, despite his constant use of the -ium spelling in all the patents[40] he filed between 1886 and 1903.[47] It has consequently been suggested that the spelling reflects an easier to pronounce word with one fewer syllable, or that the spelling on the flier was a mistake. Hall's domination of production of the metal ensured that the spelling aluminum became the standard in North America; the Webster Unabridged Dictionary of 1913, though, continued to use the -ium version. In 1926, the American Chemical Society officially decided to use aluminum in its publications; American dictionaries typically label the spelling aluminium as a British variant.
Present-day spelling Most countries spell aluminium with an i before -um. In the United States, the spelling aluminium is largely unknown, and the spelling aluminum predominates.[48] [49] The Canadian Oxford Dictionary prefers aluminum, whereas the Australian Macquarie Dictionary prefers aluminium. The International Union of Pure and Applied Chemistry (IUPAC) adopted aluminium as the standard international name for the element in 1990, but three years later recognized aluminum as an acceptable variant. Hence their periodic table includes both.[50] IUPAC officially prefers the use of aluminium in its internal publications, although several IUPAC publications use the spelling aluminum.[51]
Health concerns Despite its natural abundance, aluminium has no known function in living cells and presents some toxic effects in elevated concentrations. Its toxicity can be traced to deposition in bone and the central nervous system, which is particularly increased in patients with reduced renal function. Because aluminium competes with calcium for absorption, increased amounts of dietary aluminium may contribute to the reduced skeletal mineralization (osteopenia) observed in preterm infants and infants with growth retardation. In very high doses, aluminium can cause neurotoxicity, and is associated with altered function of the blood-brain barrier.[52] A small percentage of people are allergic to aluminium and experience contact dermatitis, digestive disorders, vomiting or other symptoms upon contact or ingestion of products containing aluminium, such as deodorants or antacids. In those without allergies, aluminium is not as toxic as heavy metals, but there is evidence of some toxicity if it is consumed in excessive amounts.[53] Although the use of aluminium cookware has not been shown to lead to aluminium toxicity in general, excessive consumption of antacids containing aluminium compounds and excessive use of aluminium-containing antiperspirants provide more significant exposure levels. Studies have shown that consumption of acidic foods or liquids with aluminium significantly increases aluminium absorption,[54] and maltol has been shown to increase the accumulation of aluminium in nervous and osseus tissue.[55] Furthermore, aluminium increases estrogen-related gene expression in human breast cancer cells cultured in the laboratory.[56] These salts' estrogen-like effects have led to their classification as a metalloestrogen.
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Aluminium Because of its potentially toxic effects, aluminium's use in some antiperspirants, dyes (such as aluminum lake), and food additives is controversial. Although there is little evidence that normal exposure to aluminium presents a risk to healthy adults,[57] several studies point to risks associated with increased exposure to the metal. Aluminium in food may be absorbed [58] more than aluminium from water. Some researchers have expressed concerns that the aluminium in antiperspirants may increase the risk of breast cancer,[59] and aluminium has controversially been implicated as a factor in Alzheimer's disease.[60] According to The Alzheimer's Society, the overwhelming medical and scientific opinion is that studies have not convincingly demonstrated a causal relationship between aluminium and Alzheimer's disease.[61] Nevertheless, some studies cite aluminium exposure as a risk factor for Alzheimer's disease, as some brain plaques have been found to contain increased levels of the metal. Research in this area has been inconclusive; aluminium accumulation may be a consequence of the disease rather than a causal agent. In any event, if there is any toxicity of aluminium, it must be via a very specific mechanism, since total human exposure to the element in the form of naturally occurring clay in soil and dust is enormously large over a lifetime.[62] [63] Scientific consensus does not yet exist about whether aluminium exposure could directly increase the risk of Alzheimer's disease.[61]
Effect on plants Aluminium is primary among the factors that reduce plant growth on acid soils. Although it is generally harmless to plant growth in pH-neutral soils, the concentration in acid soils of toxic Al3+ cations increases and disturbs root growth and function.[64] [65] [66] Most acid soils are saturated with aluminium rather than hydrogen ions. The acidity of the soil is therefore a result of hydrolysis of aluminium compounds.[67] This concept of "corrected lime potential"[68] to define the degree of base saturation in soils became the basis for procedures now used in soil testing laboratories to determine the "lime requirement" of soils.[69] Wheat's adaptation to allow aluminium tolerance is such that the aluminium induces a release of organic compounds that bind to the harmful aluminium cations. Sorghum is believed to have the same tolerance mechanism. The first gene for aluminium tolerance has been identified in wheat. It was shown that sorghum's aluminium tolerance is controlled by a single gene, as for wheat.[70] This is not the case in all plants.
See also • • • • • • •
Aluminium alloy Aluminium battery Aluminium foil Beverage can Institute for the History of Aluminium (IHA) List of countries by aluminium production Aluminium industry in Russia
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Aluminium
15
External links • • • • • •
WebElements.com – Aluminium [71] Electrolytic production [72] World production of primary aluminium, by country [73] Price history of aluminum, according to the IMF [74] History of Aluminium (from the website of the International Aluminium Institute) Emedicine - Aluminium [76]
[75]
pnb:مینمولا
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Aluminium [75] http:/ / www. world-aluminium. org/ About+ Aluminium/ Story+ of/ In+ history [76] http:/ / www. emedicine. com/ med/ topic113. htm
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Article Sources and Contributors
Article Sources and Contributors Aluminium Source: http://en.wikipedia.org/w/index.php?oldid=308237302 Contributors: (jarbarf), 050555, 111112, 56, 842U, A Brave New World, A More Perfect Onion, A.C. Norman, A.Ou, A2Kafir, ABF, AOEU, Aazn, Abrech, Accurrent, Adambro, Adashiel, Addshore, Adimovk5, AdjustShift, Aeluwas, Ahoerstemeier, Aiden Fisher, Aitias, Ajbxo, Akamad, Alansohn, Aldie, Ale jrb, Aleenf1, Alexey Feldgendler, Alexfusco5, Alexius08, Alexwcovington, AllGloryToTheHypnotoad, Allstarecho, Almazi, Alproffesor, Altenmann, Alumimum, American Eagle, Amgreen, Andesk, AndonicO, Andrevan, Andyboza, Angela, Ann Stouter, Anonymous Dissident, AnotherBrian, Ansell, Antandrus, Antonio Lopez, Anwar saadat, Arakunem, Aranherunar, Archfalhwyl, Archimerged, ArielGold, Ark2120, Army1987, Arnoutf, Artaxiad, Asder smit, Askewchan, Aslonline, Astral, At the speed of light, Atlant, Audrius u, Avala, Average Earthman, Avriette, AxSkov, AxelBoldt, Banes, Banpei, Bart Versieck, Basicdesign, Bast ailuros, Beetstra, BenFrantzDale, Beni-Sakur-Six, Benjah-bmm27, Benjiboi, Benwildeboer, BigChicken, BigHairRef, BillFlis, Billde, BillySharps, Bjarki S, Black Kite, Blainster, Blathnaid, Bletch, Bobblewik, Bobo192, Boccobrock, Boffy b, Bomac, Bongwarrior, Boothy443, Bowlhover, Brandonlovesflasks, Brandonrush, Brian0918, Brianga, Bridesmill, Brighterorange, Brion VIBBER, Britzingen, Bryan Derksen, Bubba hotep, BuickCenturyDriver, BunnyFlying, Burgercat, Bücherwürmlein, C0nanPayne, CDThieme, CMILC, CWii, CYD, Cafzal, Caltas, CambridgeBayWeather, Camden7, Camembert, Camw, Can't sleep, clown will eat me, Canadian Scouter, Canadian-Bacon, CanadianCaesar, CanadianLinuxUser, CanisRufus, Canthusus, CapitalLetterBeginning, Capricorn42, Captain panda, CaptainVindaloo, Carl Logan, Carnildo, Casliber, Casper2k3, Cassivs, Causesobad, Ccrrccrr, Celarnor, Centrx, Cfailde, Chameleon, CharlotteWebb, Chasingsol, Chatt26, ChemNerd, Chemical Heritage, Chemkid1, Chenzw, Chodges, Chris 73, Chrislk02, Chun-hian, Cireshoe, Ckatz, Clark89, Closedmouth, Cohesion, Cometstyles, CompIsMyRx, Conversion script, Corpx, CosineKitty, Coughinink, CrazyChemGuy, Cryptic C62, Curps, D, D. Recorder, D99figge, DMacks, DUBJAY04, DVD R W, Darklilac, Darrien, DarthShrine, Davehi1, Daveisrategud, Davewild, David Latapie, David.Mestel, Dawn Bard, Dbtfz, Dcljr, Dcoetzee, Ddurant, DeadEyeArrow, Decltype, Deelkar, Deglr6328, Deli nk, Delirium, Delldot, Delta G, Denni, Deor, DerHexer, Dforest, Dfrg.msc, Dhart, Diderot, Digeridouble, Dina, Discospinster, Dlohcierekim, Dmn, Donarreiskoffer, Doovinator, DoubleBlue, Doulos Christos, DrBob, DragonflySixtyseven, Drbreznjev, Dreadstar, Drutt, Dtgriscom, Duckysmokton, Dududuh, Duh duh duh, Dwmyers, Dycedarg, DylanW, Dysprosia, EL Willy, Eadthem, Eagle99999, Eaolson, East718, EdBever, Edgar181, El C, Elassint, Elektron, Element16, Eleuther, Eli the Bearded, Elkman, Emassey2005, EmilyBwater, Emperorbma, Enkauston, Enviroboy, Epbr123, Epo, Erebus555, Eric Kvaalen, Eric119, Erik Zachte, Esmehwp, Espoo, Eternal Pink, Evand, Everyking, Evil Monkey, Excirial, Exert, Eyeguy3, Ezshay, Fabiform, Fahadsadah, Faithlessthewonderboy, Faradayplank, FastMarkets, Fattyjwoods, Fayenatic london, Femto, FengRail, Ferkelparade, Fg2, Fieldday-sunday, Firebladed, FisherQueen, Flyers13, Flyguy649, FlyingToaster, Foobar, Frankenpuppy, Fraser J Allison, Freakofnurture, Freedom fighter, Freestyle-69, Friginator, Futuraind, Fvasconcellos, GPHemsley, GW Simulations, Gaff, Gail, GargoyleMT, Gary King, Gcsuchemistry, Gdo01, Gene Nygaard, Geni, Geniac, Gensanders, George The Dragon, GeorgeLouis, Georgec95, Georgewilliamherbert, Ghosts&empties, Giftlite, Gihanuk, Gilliam, Gillyweed, Giorgio51, Glenn, GlobeGores, Gnj, Gogo Dodo, GraemeL, Great Cthulhu, GreatWhiteNortherner, Gregorian, Grendelkhan, Greyhood, Grim-Gym, Grmagne, Gtstricky, Gunde123456789, Gurch, Gurubrahma, Gus Polly, Gwernol, Gyro199259, Gzuckier, Gökhan, Ha Ha, king of Switzerland, Hadal, Hadcoforyou, Haeleth, Hairy Dude, Hak-kâ-ngìn, HalfShadow, Halvorseno, Hankwang, Hannah375, Harmil, Hawk512, Hdt83, Heartache95, Hello32020, Henrik, HenryLi, Herbee, HereToHelp, Heron, HexaChord, Hgrosser, Hqb, Hraefen, Hu12, Huangcjz, Huangdi, Hut 8.5, Hydrogen Iodide, I80and, II MusLiM HyBRiD II, ITasteLikePaint, IW.HG, IWhisky, Icairns, Icek, Icepunisher, Icewedge, Igoldste, Ilurker, ImperfectlyInformed, Imroy, Inprov, InvertRect, Iridescent, Isis, Isopropyl, IstvanWolf, Iwanttodie, J.delanoy, J295, JForget, JG17, JLaTondre, JNW, JaGa, Jaan513, Jaaronan, Jacek Kendysz, JackLumber, Jackol, Jadtnr1, Jadzianna, Jaganath,
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19
Image Sources, Licenses and Contributors
Image Sources, Licenses and Contributors file:cubic-face-centered.svg Source: http://en.wikipedia.org/w/index.php?title=File:Cubic-face-centered.svg License: GNU Free Documentation License Contributors: User:Stannered file:Electron shell 013 Aluminium.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_013_Aluminium.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 file:Aluminum Metal coinless.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Aluminum_Metal_coinless.jpg License: Public Domain Contributors: User:CapitalSasha File:Loudspeaker.svg Source: http://en.wikipedia.org/w/index.php?title=File:Loudspeaker.svg License: Public Domain Contributors: Bayo, Gmaxwell, Husky, Iamunknown, Nethac DIU, Omegatron, Rocket000, 5 anonymous edits Image:Bauxite hérault.JPG Source: http://en.wikipedia.org/w/index.php?title=File:Bauxite_hérault.JPG License: unknown Contributors: saphon Image:Aluminium - world production trend.svg Source: http://en.wikipedia.org/w/index.php?title=File:Aluminium_-_world_production_trend.svg License: Public Domain Contributors: User:Leyo Image:Aluminium output2.PNG Source: http://en.wikipedia.org/w/index.php?title=File:Aluminium_output2.PNG License: Public Domain Contributors: User:Anwar_saadat/bubble_maps_(FAQ) File:41 ALU Recycling Code.svg Source: http://en.wikipedia.org/w/index.php?title=File:41_ALU_Recycling_Code.svg License: GNU Free Documentation License Contributors: Karl A Randall / User: k4rlR Image:aluminumfoil.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Aluminumfoil.jpg License: Public Domain Contributors: Kerkyra, 1 anonymous edits Image:Austin A40 Roadster ca 1951.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Austin_A40_Roadster_ca_1951.jpg License: unknown Contributors: User:Charles01 Image:AluminumSlab.JPG Source: http://en.wikipedia.org/w/index.php?title=File:AluminumSlab.JPG License: Public Domain Contributors: User:ALIquotob Image:Aluminium foam.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Aluminium_foam.jpg License: unknown Contributors: User:Stehfun Image:Eros-piccadilly-circus.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Eros-piccadilly-circus.jpg License: GNU Free Documentation License Contributors: Bhoeble, Doruk Salancı, G.dallorto, Hanzo86, Jastrow, Joanjoc, Justinc, Korrigan, Montrealais, Pseudomoi, Saperaud, Wst, 3 anonymous edits
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
20
Silicon
1
Silicon aluminium ← silicon → phosphorusC ↑ Si ↓ Ge
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HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
14Si Periodic table
Appearance crystalline, reflective with bluish-tinged faces
Broken
silicon
ingot
General
Name,
symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
Silicon
2 at T/K
1908
2102
2339
2636
3021
3537
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1577.1 kJ·mol−1 3rd: 3231.6 kJ·mol−1Atomic radiusCovalent radiusVan der Waals radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusPoisson ratioMohs hardnessCAS registry numberBand gap energy at 300 K Most stable isotopes Main article: Isotopes of silicon iso
N.A.
half-life
28
92.23%
28
29
4.67%
29
30
3.1%
30
32
syn
170 y
Si Si Si Si
DM
DE (MeV)
DP
Si is stable with 14 neutron Si is stable with 15 neutron Si is stable with 16 neutron β−
13.020
32
P
silicon, Si, 14 metalloid14, 3, p28.0855(3) g·mol−1 [Ne] 3s2 3p2 2, 8, 4 (Image) solid 2.3290 g·cm−3 2.57 g·cm−3 1687 K,1414 °C,2577 °F 3538 K,3265 °C,5909 °F 50.21 kJ·mol−1 359 kJ·mol−1 (25 °C) 19.789 J·mol−1·K−14, 3 , 2 , 1[1] -1, -2, -3, -4 (amphoteric oxide) 1.90 (Pauling scale) 1st: 786.5 kJ·mol−1111 pm111 pm 210 pm diamond cubic diamagnetic[2] (20 °C) 103 [3] Ω·m (300 K) 149 W·m−1·K−1 (25 °C) 2.6 µm·m−1·K−1 (20 °C) 8433 m/s 185[3] GPa 52[3] GPa 100 GPa 0.28[3] 7 7440-21-3 1.12 eV Silicon (pronounced /ˈsɪlɨkən/ or English pronunciation: /ˈsɪlɨkɒn/, Latin: silicium) is the most common metalloid. It is a chemical element, which has the symbol Si and atomic number 14. The atomic mass is 28.0855. A tetravalent metalloid, silicon is less reactive than its chemical analog carbon. As the eighth most common element in the universe by mass, silicon very rarely occurs as the pure free element in nature, but is more widely distributed in dusts, planetoids and planets as various forms of silicon dioxide (silica) or silicates. On Earth, silicon is the second most abundant element (after oxygen) in the crust,[4] making up 25.7% of the crust by mass. Silicon has many industrial uses. It is the principal component of most semiconductor devices, most importantly integrated circuits or microchips. Silicon is widely used in semiconductors because it remains a semiconductor at higher temperatures than the semiconductor germanium and because its native oxide is easily grown in a furnace and forms a better semiconductor/dielectric interface than any other material. In the form of silica and silicates, silicon forms useful glasses, cements, and ceramics. It is also a constituent of silicones, a class-name for various synthetic plastic substances made of silicon, oxygen, carbon and hydrogen, often confused with silicon itself. Silicon is an essential element in biology, although only tiny traces of it appear to be required by animals.[5] It is much more important to the metabolism of plants, particularly many grasses, and silicic acid (a type of silica) forms the basis of the striking array of protective shells of the microscopic diatoms.
Silicon
Notable characteristics The outer electron orbitals (half filled subshell holding up to eight electrons) have the same structure as in carbon and the two elements are sometimes similar chemically. Even though it is a relatively inert element, silicon still reacts with halogens and dilute alkalis, but most acids (except for some hyper-reactive combinations of nitric acid and hydrofluoric acid) do not affect it. Having four bonding electrons however gives it, like carbon, many opportunities to combine with other elements or compounds under the right circumstances. Both silicon and (in certain aspects) carbon are semiconductors, readily either donating or sharing their four outer electrons allowing many different forms of chemical bonding. Pure silicon has a negative temperature coefficient of resistance, since the number of free charge carriers increases with temperature. The electrical resistance of single crystal silicon significantly changes under the application of mechanical stress due to the piezoresistive effect. In its crystalline form, pure silicon has a gray color and a metallic luster. It is similar to glass in that it is rather strong, very brittle, and prone to chipping.
History Silicon was first identified by Antoine Lavoisier in 1787 (as a component of the Latin silex, silicis for flint, flints), and was later mistaken by Humphry Davy in 1800 for a compound. In 1811 Gay-Lussac and Thénard probably prepared impure amorphous silicon through the heating of potassium with silicon tetrafluoride. In 1824, Berzelius prepared amorphous silicon using approximately the same method as Lussac. Berzelius also purified the product by repeatedly washing it.[6]
Occurrence Measured by mass, silicon makes up 25.7% of the Earth's crust and is the second most abundant element in the crust, after oxygen. Pure silicon crystals are very rarely found in nature; they can be found as inclusions with gold and in volcanic exhalations. Silicon is usually found in the form of silicon dioxide (also known as quartz), and other more complex silicate minerals. Silica occurs in minerals consisting of (practically) pure silicon dioxide in different crystalline forms. Amethyst, agate, quartz, rock crystal, chalcedony, flint, jasper, and opal are some of the forms in which silicon dioxide appears. Biogenic silica occurs in the form of diatoms, radiolaria and siliceous sponges. Silicon also occurs as silicate minerals (various minerals containing silicon, oxygen and one or another metal), for example the feldspar group. These minerals occur in clay, sand and various types of rock such as granite and sandstone. Feldspar, pyroxene, amphibole, and mica are a few of the many common silicate mineral groups. Silicon is a principal component of many meteorites, and also is a component of obsidian and tektites, which are natural forms of glass. See also Category:Silicate minerals
3
Silicon
4
Isotopes Silicon has numerous known isotopes, with mass numbers ranging from 22 to 44. 28Si (the most abundant isotope, at 92.23%), 29Si (4.67%), and 30Si (3.1%) are stable; 32Si is a radioactive isotope produced by cosmic ray spallation of argon. Its half-life has been determined to be approximately 170 years (0.21 MeV), and it decays by beta - emission to 32 P (which has a 14.28 day half-life )[7] and then to 32S.
Production Silicon is commercially prepared by the reaction of high-purity silica with wood, charcoal, and coal, in an electric arc furnace using carbon electrodes. At temperatures over 1900 °C (3450 °F), the carbon reduces the silica to silicon according to the chemical equations: SiO2 + C → Si + CO2 SiO2 + 2 C → Si + 2 CO Liquid silicon collects in the bottom of the furnace, and is then drained and cooled. The silicon produced via this process is called metallurgical grade silicon and is at least 98% pure. Using this method, silicon carbide (SiC) may form. However, provided the concentration of SiO2 is kept high, the silicon carbide can be eliminated: 2 SiC + SiO2 → 3 Si + 2 CO In September 2008, metallurgical grade silicon cost about USD 1.45 per pound ($3.20/kg),[8] up from $0.77 per pound ($1.70/kg) in 2005.[9] Recently, it has been reported that pure silicon can be extracted directly from solid silica by molten salt electrolysis. This new method, known as the FFC Cambridge Process, has the potential to directly produce solar grade silicon without any CO2 emission and at much lower energy consumption.[10] [11] [12]
Crystallization Silicon, like carbon and other group IV elements form face-centered diamond cubic crystal structure. Silicon, in particular, forms a face-centered cubic structure with a lattice spacing of 5.430710 Å (0.5430710 nm).[13] The majority of silicon crystals grown for device production are produced by the Czochralski process, (CZ-Si) since it is the cheapest method available and it is capable of producing large size crystals. However, silicon single-crystals grown by the Czochralski method contain impurities since the crucible which contains the melt dissolves. For certain electronic devices, Diamond Cubic Crystal Structure, particularly those required for high power applications, Silicon unit cell silicon grown by the Czochralski method is not pure enough. For these applications, float-zone silicon (FZ-Si) can be used instead. It is worth mentioning though, in contrast with CZ-Si method in which the seed is dipped into the silicon melt and the growing crystal is pulled upward, the thin seed crystal in the FZ-Si method sustains the growing crystal as well as the polysilicon
Silicon
5
rod from the bottom. As a result, it is difficult to grow large size crystals using the float-zone method. Today, all the dislocation-free silicon crystals used in semiconductor industry with diameter 300 mm or larger are grown by the Czochralski method with purity level significantly improved.
Purification The use of silicon in semiconductor devices demands a much greater purity than afforded by metallurgical grade silicon. Historically, a number of methods have been used to produce high-purity silicon.
Physical methods Early silicon purification techniques were based on the fact that if silicon is melted and re-solidified, the last parts of the mass to solidify contain most of the impurities. The earliest method of silicon purification, first described in 1919 and used on a limited basis to make radar components during World War II, involved crushing metallurgical grade silicon and then partially dissolving the silicon powder in an acid. When crushed, the silicon cracked so that the weaker impurity-rich regions were on the outside of the resulting grains of silicon. As a result, the impurity-rich silicon was the first to be dissolved when treated with acid, leaving behind a more pure product. In zone melting, also called zone refining, the first silicon purification method to be widely used industrially, rods of metallurgical grade silicon are heated to melt at one end. Then, the heater is slowly moved down the length of the rod, keeping Silicon ingot a small length of the rod molten as the silicon cools and re-solidifies behind it. Since most impurities tend to remain in the molten region rather than re-solidify, when the process is complete, most of the impurities in the rod will have been moved into the end that was the last to be melted. This end is then cut off and discarded, and the process repeated if a still higher purity is desired.
Chemical methods Today, silicon is purified by converting it to a silicon compound that can be more easily purified by distillation than in its original state, and then converting that silicon compound back into pure silicon. Trichlorosilane is the silicon compound most commonly used as the intermediate, although silicon tetrachloride and silane are also used. When these gases are blown over silicon at high temperature, they decompose to high-purity silicon. At one time, DuPont produced ultra-pure silicon by reacting silicon tetrachloride with high-purity zinc vapors at 950 °C, producing silicon: SiCl4 + 2 Zn → Si + 2 ZnCl2
Silicon
6
However, this technique was plagued with practical problems (such as the zinc chloride byproduct solidifying and clogging lines) and was eventually abandoned in favor of the Siemens process. In the Siemens process, high-purity silicon rods are exposed to trichlorosilane at 1150 °C. The trichlorosilane gas decomposes and deposits additional silicon onto the rods, enlarging them: 2 HSiCl3 → Si + 2 HCl + SiCl4 Silicon produced from this and similar processes is called polycrystalline silicon. Polycrystalline silicon typically has impurity levels of less than 10−9. In 2006 REC announced construction of a plant based on fluidized bed technology using silane:[14]
A polycrystalline silicon rod made by the Siemens process
3 SiCl + Si + 2 H → 4 HSiCl 4
2
3
4 HSiCl3 → 3 SiCl4 + SiH4 SiH4 → Si + 2 H2
Different forms of silicon
Nanocrystalline silicon powder Silicon powder
Silicon wafer with mirror finish (NASA)
One can notice the colour change in silicon nanopowder. This is caused by the quantum effects which occur in particles of nanometric dimensions. See also Potential well, Quantum dot, and Nanoparticle.
Silicon
Compounds Silicon forms binary compounds called silicides with many metallic elements whose properties range from reactive compounds, e.g. magnesium silicide, Mg2Si through high melting refractory compounds such as molybdenum disilicide, MoSi2.[15] Silicon carbide, SiC (carborundum) is a hard, high melting solid and a well known PDMS – a silicone abrasive. Silane, SiH4, is a pyrophoric gas with a similar tetrahedral compound structure to methane, CH4. Additionally there is a range of catenated silicon hydrides that form a homologous series of compounds, SinH2n+2 where n = 2-8 (analogous to the alkanes).[16] These are all readily hydrolysed and are thermally unstable, particularly the heavier members.[16] Disilenes contain a silicon-silicon double bond (analogous to the alkenes) and are generally highly reactive requiring large substituent groups to stabilise them.[17] A disilyne with a silicon-silicon triple bond was first isolated in 2004; although as the compound is non-linear, the bonding is dissimilar to that in alkynes.[18] Tetrahalides, SiX4, are formed with all of the halogens.[15] Silicon tetrachloride, for example, readily reacts with water; unlike its carbon analogue, carbon tetrachloride.[16] Silicon dihalides are formed by the high temperature reaction of tetrahalides and silicon; with a structure analogous to a carbene they are reactive compounds.[16] Silicon difluoride condenses to form a polymeric compound, (SiF2)n.[16] Silicon dioxide is a high melting solid with a number of different crystal forms; the most familiar of which is the mineral quartz.[15] In quartz each silicon atom is surrounded by four oxygen atoms that bridge to other silicon atoms to form a three dimensional lattice.[15] Silica is soluble in water at high temperatures forming monosilicic acid, (Si(OH)4)[16] and this property is used in the manufacture of quartz crystals used in electronics.[15] Under the right conditions monosilicic acid readily polymerises to form more complex silicic acids, ranging from the simplest condensate, disilicic acid (H6Si2O7) to linear, ribbon, layer and lattice structures which form the basis of the many different silicate minerals.[16] Silicates are also important constituents of concretes.[15] With oxides of other elements the high temperature reaction of silicon dioxide can give a wide range of glasses with various properties.[16] Examples include soda lime glass, borosilicate glass and lead crystal glass. Silicon sulfide, SiS2 is a polymeric solid (unlike its carbon analogue the liquid CS2).[15] Silicon forms a nitride, Si3N4 which is a ceramic.[15] Silatranes, a group of tricyclic compounds containing five-coordinate silicon, may have physiological properties.[19] Many transition metal complexes containing a metal-silicon bond are now known, which include complexes containing SiHnX3−n ligands, SiX3 ligands, and Si(OR)3 ligands.[19] Silicones are large group of polymeric compounds with an (Si-O-Si) backbone. An example is the silicone oil PDMS (polydimethylsiloxane).[15] These polymers can be crosslinked to produce resins and elastomers.[15] Many organosilicon compounds are known which contain a silicon-carbon single bond. Many of these are based on a central tetrahedral silicon atom, and some are optically active when central chirality exists. Long chain polymers containing a silicon backbone are known, such as polydimethysilylene (SiMe2)n.[20] Polycarbosilane, [(SiMe2)2CH2]n with a backbone containing a repeating -Si-Si-C unit, is a precursor in the production of silicon carbide fibres.[20]
7
Silicon
Applications As the second most abundant element in the earth's crust, silicon is vital to the construction industry as a principal constituent of natural stone, glass, concrete and cement. Silicon's greatest impact on the modern world's economy and lifestyle has resulted from silicon wafers used as substrates in the manufacture of discrete electronic devices such as power transistors, and in the development of integrated circuits such as computer chips.
Alloys The largest application of metallurgical grade silicon, representing about 55% of the world consumption, is in the manufacture of aluminium-silicon alloys to produce cast parts, mainly for the automotive industry. Silicon is an important constituent of electrical steel, modifying its resistivity and ferromagnetic properties. Silicon is added to molten cast iron as ferrosilicon or silicocalcium alloys to improve its performance in casting thin sections, and to prevent the formation of cementite at the surface.
In electronic applications Pure silicon is used to produce ultra-pure silicon wafers used in the semiconductor industry, in electronics and in photovoltaic applications. Ultra-pure silicon can be doped with other elements to adjust its electrical response by controlling the number and charge (positive or negative) of current carriers. Such control is necessary for transistors, solar cells, integrated circuits, microprocessors, semiconductor detectors and other semiconductor devices which are used in electronics and other high-tech applications. In silicon photonics, it can be used as a continuous wave Raman laser medium to produce coherent light, though it is ineffective as a light source. Hydrogenated amorphous silicon is used in the production of low-cost, large-area electronics in applications such as LCDs, and of large-area, low-cost thin-film solar cells.
Silicones The second largest application of silicon (about 40% of world consumption) is as a raw material in the production of silicones, compounds containing silicon-oxygen and silicon-carbon bonds that have the capability to act as bonding intermediates between glass and organic compounds, and to form polymers with useful properties such as impermeability to water, flexibility and resistance to chemical attack. Silicones are used in waterproofing treatments, molding compounds and mold-release agents, mechanical seals, high temperature greases and waxes, caulking compounds and even in applications as diverse as breast implants, contact lenses, explosives and pyrotechnics.[21] • Construction: Silicon dioxide or silica in the form of sand and clay is an important ingredient of concrete and brick and is also used to produce Portland cement. • Pottery/Enamel is a refractory material used in high-temperature material production and its silicates are used in making enamels and pottery. • Glass: Silica from sand is a principal component of glass. Glass can be made into a great variety of shapes and with many different physical properties. Silica is used as a base material to make window glass, containers, insulators, and many other useful objects. • Abrasives: Silicon carbide is one of the most important abrasives. • Silly Putty was originally made by adding boric acid to silicone oil.[22] See also Category:Silicon compounds
8
Silicon
In popular culture Because silicon is an important element in semiconductors and high-tech devices, the high-tech region of Silicon Valley, California is named after this element. Other geographic locations with connections to the industry have since characterized themselves as Siliconia as well, for example Silicon Forest in Oregon, Silicon Saxony in Germany, and Silicon Border in Mexicali.
See also • • • • • • • •
Black silicon Covalent superconductors Crystalline silicon Electronics List of silicon producers Printed silicon electronics RCA clean Silicone
• Wafer (electronics)
External links • Mineral.Galleries.com – Silicon [23] • WebElements.com – Silicon [24]
References [1] R. S. Ram et al. "Fourier Transform Emission Spectroscopy of the A2D–X2P Transition of SiH and SiD" J. Mol. Spectr. 190, 341–352 (1998) (http:/ / bernath. uwaterloo. ca/ media/ 184. pdf) [2] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [3] http:/ / www. ioffe. ru/ SVA/ NSM/ Semicond/ Si [4] " The periodic table (http:/ / www. webelements. com/ )". webelements.com. . Retrieved 2008-02-20. [5] Nielsen, Forrest H. (1984). "Ultratrace Elements in Nutrition". Annual Review of Nutrition 4: 21–41. doi: 10.1146/annurev.nu.04.070184.000321 (http:/ / dx. doi. org/ 10. 1146/ annurev. nu. 04. 070184. 000321). [6] Weeks, Mary Elvira (1932). "The discovery of the elements: XII. Other elements isolated with the aid of potassium and sodium: beryllium, boron, silicon, and aluminum". Journal of Chemical Education: 1386–1412. [7] " Phosphorus - 32 (http:/ / sciencegateway. org/ isotope/ phosp32. html)". sciencegateway.org. . Retrieved 2008-02-20. [8] " Metallurgical silicon could become a rare commodity – just how quickly that happens depends to a certain extent on the current financial crisis (http:/ / www. photon-magazine. com/ news_archiv/ details. aspx?cat=News_PI& sub=worldwide& pub=4& parent=1555)". Photon International. . Retrieved 2009-03-04. [9] " Silicon (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ silicon/ silicmcs06. pdf)". usgs.gov. . Retrieved 2008-02-20. [10] " New silicon production method with no carbon reductant (http:/ / www. acr. net. au/ ~coastwatchers/ charcoalition/ noreductant. html)". . Referering to Chen, George Zheng; Fray, Derek J.; Farthing, Tom W. (2000). "Direct electrochemical reduction of titanium dioxide to titanium in molten calcium chloride". Nature 407 (6802): 361–364. doi: 10.1038/35030069 (http:/ / dx. doi. org/ 10. 1038/ 35030069). [11] Xianbo Jin, Pei Gao, Dihua Wang, Xiaohong Hu, George Z. Chen (2004). "Electrochemical Preparation of Silicon and Its Alloys from Solid Oxides in Molten Calcium Chloride". Angew. Chem. Int. Ed. 43: 733–736. doi: 10.1002/anie.200352786 (http:/ / dx. doi. org/ 10. 1002/ anie. 200352786). [12] Nohira, Toshiyuki; Yasuda, Kouji; Ito, Yasuhiko (2003). "Pinpoint and bulk electrochemical reduction of insulating silicon dioxide to silicon". Nature Materials 2: 397–401. doi: 10.1038/nmat900 (http:/ / dx. doi. org/ 10. 1038/ nmat900).
9
Silicon [13] O'Mara, William C. (1990). Handbook of Semiconductor Silicon Technology (http:/ / books. google. com/ books?id=COcVgAtqeKkC& pg=PA351& dq=Czochralski+ Silicon+ Crystal+ Face+ Cubic& lr=& as_brr=3& sig=ht-dgSy1lzBMYC7IXPp9W5QBqYo). William Andrew Inc.. p. 349-352. ISBN 0815512376. . Retrieved 2008-02-24. [14] " Analyst_silicon_field_trip_March_28,_2007 (http:/ / hugin. info/ 136555/ R/ 1115224/ 203491. pdf)". hugin.info. . Retrieved 2008-02-20. [15] Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.). Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4. [16] Holleman, A. F.; Wiberg, E.; Wiberg, N. (2001). Inorganic Chemistry, 1st Edition. Academic Press. ISBN 0123526515. [17] (Multiply Bonded Main Group Metals and Metalloids: Multiple Bonded Main Group Metals and Metalloids, F. G. Stone, Robert West, Academic Press, 1996, ISBN 0120311399) [18] A Stable Compound Containing a Silicon-Silicon Triple Bond, Akira Sekiguchi, Rei Kinjo, Masaaki Ichinohe, Science 17 September 2004:Vol. 305. no. 5691, pp. 1755 - 1757 doi: 10.1126/science.1102209 (http:/ / dx. doi. org/ 10. 1126/ science. 1102209) [19] Lickiss, Paul D. (1994). Inorganic Compounds of Silicon, in Encyclopedia of Inorganic Chemistry. John Wiley & Sons. pp. 3770 - 3805. ISBN 0471936200. [20] Mark, James. E (2005). Inorganic polymers. Oxford University Press. pp. 200-245. ISBN 0195131193. [21] Koch, E.C.; Clement, D.. Special Materials in Pyrotechnics: VI. Silicon - An Old Fuel with New Perspectives (http:/ / www3. interscience. wiley. com/ cgi-bin/ abstract/ 114279686/ ABSTRACT). . [22] Walsh, Tim (2005). " Silly Putty (http:/ / books. google. de/ books?id=jftapGDTmYUC& pg=PA90)". Timeless toys: classic toys and the playmakers who created them. Andrews McMeel Publishing. ISBN 9780740755712. . [23] http:/ / mineral. galleries. com/ minerals/ elements/ silicon/ silicon. htm [24] http:/ / www. webelements. com/ webelements/ elements/ text/ Si/ key. html
10
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Article Sources and Contributors Silicon Source: http://en.wikipedia.org/w/index.php?oldid=308114415 Contributors: 03272, 12dstring, 2D, 45Factoid44, AB, ABF, Abby, Abrech, Addshore, Adrian.benko, Ahoerstemeier, Alansohn, Aldaron, AlexOvShaolin, AlexiusHoratius, AlistairMcMillan, Allison Connors, Allstarecho, Anclation, Andfarm, AndonicO, Andres, Andrew Nutter, Andrewmc123, Animum, Ann Stouter, Antandrus, Arakunem, Archimerged, Arthena, Astrochemist, Atlant, Aubriexoxo, AuburnPilot, AxelBoldt, Axiosaurus, Bantman, Bartholomeyczik, Barticus88, Bcasterline, Beetstra, Benbest, Bender235, Biskuvi, Bjweeks, Blanchardb, Bloodofjing, BlueEarth, Bobo192, Bomac, Brian0918, Brockert, Buster79, Bwiki, CMBJ, CSWarren, CWii, CYD, Caiaffa, Call me Bubba, Can't sleep, clown will eat me, CanisRufus, Canterbury Tail, Capricorn42, Captain panda, Card, CardinalDan, Carnildo, Cbaker9552, Cflm001, Chairboy, Chaos syndrome, CharlotteWebb, Ciaccona, Ck lostsword, Coemgenus, Colboi, Cometstyles, Communisthamster, Condem, Conversion script, Coppertwig, CrazyChemGuy, Cremepuff222, Crystal whacker, Ctjf83, Cyhatch, Cyp, DA3N, DJ Clayworth, Daano15, Dali, Daniel 1992, Daniel5127, Danski14, Dar-Ape, Darrien, Darth Panda, Davewild, David Latapie, David0811, Deglr6328, Deli nk, Delldot, Delphonic, Delta G, Denelson83, DerHexer, Dgrant, Dgwohu, Dichrra, Discospinster, Dismas, Djtrimz, DmitTrix, Donarreiskoffer, Doodle77, Double Vigie, Doulos Christos, Download, Dr bab, DrBob, Draeco, Droidus, Dspark76, Dspradau, Dwmyers, Dysepsion, Dzukman2000, E dog95, ESkog, Edgar181, El C, Eldin raigmore, Elliskev, Emperorbma, Enricoros, Enviroboy, Epbr123, Eric119, Erik Zachte, Erik9, Eskimospy, Eudialytos, EvocativeIntrigue, FSIM, FT2, Fabartus, FastLizard4, Fbriere, Femto, Fieldday-sunday, Flewis, Fonzy, Frankenpuppy, Galoubet, Gene Nygaard, Giftlite, Gimme danger, Glacier Wolf, Glengarry, Gman124, Golbez, Goron1130, Gracefool, GregorB, Grendelkhan, Gurch, Gwernol, Gökhan, Haasl, Hadal, Hak-kâ-ngìn, Hankwang, Hastings007, Hayabusa future, Hdt83, Heimstern, Helikophis, HenryLi, HereToHelp, Heroismic, Heron, HexaChord, Hobartimus, Hunt ice, Hut 8.5, Hyperdeath, Icairns, Igoldste, Imaginateca, Instinct, Inter, Iridescent, Isfisk, It.franciscus, ItaniuMatrix, Iustinus, J.delanoy, JForget, Jabam, JackSchmidt, Jamix, Jaraalbe, Jaxl, JayZ, Jclemens, Jeff G., Jeronimo, Jh51681, Jobarts, John, John254, JorisvS, Jose77, Josh Grosse, Jotomicron, Juliancolton, JustinWick, Kajerm, Kariboo, Karl-Henner, Karlhahn, Kchishol1970, Keith Lehwald, Kirsten07734, Knutux, Kojiki1976, Kotra, Kr5t, Kraxler, Kristen Eriksen, KristianMolhave, Ktsquare, Kuru, Kwamikagami, LA2, Lando5, LeaveSleaves, Lethalgeek, Linnell, Little Mountain 5, Logical2u, Looxix, LuigiManiac, Lupin, Luzheng, Lvl, MBisanz, MC Steel, MZMcBride, Mac, Makemi, Malinaccier, Mandarax, Marauder40, Marek69, Master of Pies, Materialscientist, Matt Britt, Mav, Mcorazao, MedievalFreak, Megan1967, Melah Hashamaim, Merope, Michael.j.sykora, Miciah, Miquonranger03, Miranda, Miss Madeline, Modulatum, Moriori, Mr. Qwert, Ms2ger, Murphyjd, Musser, Mxn, Myleslong, NJPharris, Naaa127, Nakon, NawlinWiki, Ndittert, Nergaal, Neverquick, NewEnglandYankee, Nielmo, NightFalcon90909, Nihiltres, Nlu, No Guru, Nuttycoconut, Oda Mari, Oldboltonian, Olorin28, Omegatron, Omicronpersei8, Onevalefan, Opelio, Ortolan88, Ossmann, Oxymoron83, P.B. 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11
Phosphorus
1
Phosphorus silicon ← phosphorus → sulfurN ↑ P ↓ As
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HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
15P Periodic table
Appearance colorless, waxy white, yellow, scarlet, red, violet, black
waxy white (yellow cut), red, violet and black phoshorus General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties Density (near r.t.) Melting pointSublimation pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure (white) P/Pa
1
10
100
1k
10 k
100 k
at T/K
279
307
342
388
453
549
Vapor pressure (red, bp. 431 °C) P/Pa
1
10
100
1k
10 k
100 k
at T/K
455
489
529
576
635
704
Phosphorus
2
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1907 kJ·mol−1 3rd: 2914.1 kJ·mol−1Covalent radiusVan der Waals radius Miscellaneous Magnetic orderingThermal conductivityBulk modulusCAS registry number Most stable isotopes Main article: Isotopes of phosphorus iso
N.A.
half-life
DM
DE (MeV)
DP
31
100%
31
32
syn
14.28 d
β−
1.709
32
33
syn
25.3 d
β−
0.249
33
P P P
P is stable with 16 neutron S S
phosphorus, P, 15 nonmetal15, 3, p30.973762(2) g·mol−1 [Ne] 3s2 3p3 2, 8, 5 (Image) (white) 1.823, (red) ≈ 2.2 – 2.34, (violet) 2.36, (black) 2.69 g·cm−3 (white) 44.2 °C, (black) 610 °C (red) ≈ 416 – 590 °C, (violet) 620 °C (white) 280.5 °C (white) 0.66 kJ·mol−1 (white) 12.4 kJ·mol−1 (25 °C) (white) 23.824 J·mol−1·K−15, 4, 3, 2[1] , 1 [2] , -1, -2, -3 (mildly acidic oxide) 2.19 (Pauling scale) 1st: 1011.8 kJ·mol−1107±3 pm 180 pm (white,red,violet,black) diamagnetic[3] (300 K) (white) 0.236, (black) 12.1 W·m−1·K−1 (white) 5, (red) 11 GPa 7723-14-0 Phosphorus (pronounced /ˈfɒsfərəs/) is the chemical element that has the symbol P and atomic number 15. A multivalent nonmetal of the nitrogen group, phosphorus is commonly found in inorganic phosphate rocks. Elemental phosphorus exists in two major forms - white phosphorus and red phosphorus. Although the term "phosphorescence", meaning glow after illumination, derives from phosphorus, glow of phosphorus originates from oxidation of the white (but not red) phosphorus and should be called chemiluminescence. Due to its high reactivity, phosphorus is never found as a free element in nature on Earth. The first form of phosphorus to be discovered (white phosphorus, discovered in 1669) emits a faint glow upon exposure to oxygen — hence its name given from Greek mythology, Φωσφόρος meaning "light-bearer" (Latin Lucifer), referring to the "Morning Star", the planet Venus. Phosphorus is a component of DNA, RNA, ATP, and also the phospholipids which form all cell membranes. It is thus an essential element for all living cells. The most important commercial use of phosphorus-based chemicals is the production of fertilizers. Phosphorus compounds are also widely used in explosives, nerve agents, friction matches, fireworks, pesticides, toothpaste and detergents.
Physical properties Glow from white phosphorus Phosphorus was discovered by German alchemist Hennig Brand in 1674 or 1675. Working in Hamburg, Brand attempted to distil some kind of "life essence" from his urine, and in the process produced a white material that glowed in the dark.[4] The phosphorus had in fact been produced from inorganic phosphate, which is a significant component of dissolved urine solids. White phosphorus is highly reactive and gives off a faint greenish glow upon uniting with oxygen. The glow observed by Brand was actually caused by the very slow burning of the phosphorus, but as he neither saw flame nor felt any heat he did not
Phosphorus
3
recognize it as burning. It was known from early times that the glow would persist for a time in a stoppered jar but then cease. Robert Boyle in the 1680s ascribed it to "debilitation" of the air; in fact, it is oxygen being consumed. By the 18th century, it was known that in pure oxygen, phosphorus does not glow at all;[5] there is only a range of partial pressure at which it does. Heat can be applied to drive the reaction at higher pressures.[6] In 1974, the glow was explained by R. J. van Zee and A. U. Khan.[7] A reaction with oxygen takes place at the surface of the solid (or liquid) phosphorus, forming the short-lived molecules HPO and P2O2 that both emit visible light. The reaction is slow and only very little of the intermediates are required to produce the luminescence, hence the extended time the glow continues in a stoppered jar. Although the term phosphorescence is derived from phosphorus, the reaction which gives phosphorus its glow is properly called chemiluminescence (glowing due to a cold chemical reaction), not phosphorescence (re-emitting light that previously fell onto a substance and excited it). Phosphorescence is the slow decay of a metastable electronic state to a lower energy state through emission of light. The decay is slow because the transition from the excited to the lower state requires a spin flip, making it classically forbidden. Often it involves a transition from an excited triplet state to a singlet ground state. The metastable excited state may have been populated by thermal excitations or some light source. Since phosphorescence is slow, it persists for some time after the exciting source is removed. In contrast, chemiluminescence occurs when the product molecules of a chemical reaction (HPO and P2O2 in this case) leave the reaction in an electronically excited state. These excited molecules then release their excess energy in the form of light. The frequency (colour) of the light emitted is proportional to the energy difference of the two electronic states involved.[8]
Allotropes Phosphorus has several forms (allotropes) which have strikingly different properties.[9] The two most common allotropes are white phosphorus and red phosphorus. Red phosphorus is an intermediate phase between white and violet phosphorus. Another form, scarlet phosphorus, is obtained by allowing a solution of white phosphorus in carbon disulfide to evaporate in sunlight. Black phosphorus is obtained by heating white phosphorus under high pressures (about P4 molecule 12,000 atmospheres). In appearance, properties, and structure, it resembles graphite, being black and flaky, a conductor of electricity, and has puckered sheets of linked atoms. Another allotrope is diphosphorus; it contains a phosphorus dimer as a structural unit and is highly reactive.[]
Phosphorus
4
White phosphorus has two forms, low-temperature β form and high-temperature α form. They both contain a phosphorus P4 tetrahedron as a structural unit, in which each atom is bound to the other three atoms by a single bond. This P4 tetrahedron is also present in liquid and gaseous phosphorus up to the temperature of 800 °C when it starts decomposing to P2 molecules.[10] White phosphorus is the least stable, the most reactive, more volatile, less P4O10 molecule dense, and more toxic than the other allotropes. The toxicity of white phosphorus led to its discontinued use in matches. White phosphorus is thermodynamically unstable at normal condition and will gradually change to red phosphorus. This transformation, which is accelerated by light and heat, makes white phosphorus almost always contain some red phosphorus and therefore appear yellow. For this reason, it is also called yellow phosphorus. It glows greenish in the dark (when exposed to oxygen), is highly flammable and pyrophoric (self-igniting) upon contact with air as well as toxic (causing severe liver damage on ingestion). Because of pyrophoricity, white phosphorus is used as an additive in napalm. The odour of combustion of this form has a characteristic garlic smell, and samples are commonly coated with white "(di)phosphorus pentoxide", which consists of P4O10 tetrahedra with oxygen inserted between the phosphorus atoms and at their vertices. White phosphorus is insoluble in water but soluble in carbon disulfide.[11] The white allotrope can be produced using several different methods. In one process, calcium phosphate, which is derived from phosphate rock, is heated in an electric or fuel-fired furnace in the presence of carbon and silica.[12] Elemental phosphorus is then liberated as a vapour and can be collected under phosphoric acid. This process is similar to the first synthesis of phosphorus from calcium phosphate in urine. In the red phosphorus, one of the P bonds is broken, 4
and one additional bond is formed with a neighbouring tetrahedron resulting in a more chain-like structure. Red phosphorus may be formed by heating white phosphorus to 250 °C (482 °F) or by exposing white phosphorus to sunlight.[13] Phosphorus after this Crystal structure of red phosphorus treatment exists as an amorphous network of atoms which reduces strain and gives greater stability; further heating results in the red phosphorus becoming crystalline. Therefore red phosphorus is not a certain allotrope, but rather an intermediate phase between the white and violet phosphorus, and most of its properties have a range of values. Red phosphorus does not catch fire in air at temperatures below 260 °C, whereas white phosphorus ignites at about 30 °C.[14] Violet phosphorus is a thermodynamic stable form of phosphorus which can be produced by day-long temper of red phosphorus above 550 °C. In 1865, Hittorf discovered that when phosphorus was recrystallized from molten lead, a red/purple form is obtained. Therefore this form is sometimes known as "Hittorf's phosphorus" (or violet or α-metallic phosphorus).[]
Phosphorus
5
Black phosphorus is the least reactive allotrope and the thermodynamic stable form below 550 °C. It is also known as β-metallic phosphorus and has a structure somewhat resembling that of graphite.[15] [16] High pressures are usually required to produce black phosphorus, but it can also be produced at ambient conditions using metal salts as catalysts.[17] The diphosphorus allotrope, P2, is stable only at high Crystal structure of black temperatures. The dimeric unit contains a triple bond and is phosphorus analogous to N2. The diphosphorus allotrope (P2) can be obtained normally only under extreme conditions (for example, from P4 at 1100 kelvin). Nevertheless, some advancements were obtained in generating the diatomic molecule in homogeneous solution, under normal conditions with the use by some transitional metal complexes (based on, for example, tungsten and niobium).[18]
Properties of some allotropes of phosphorus Form Symmetry
white(α) Body-centred cubic
Pearson symbol
white(β)
violet
black
Triclinic
Monoclinic
Orthorhombic
aP24
mP84
oS8
Space group
I-43m
P-1 No.2
P2/c No.13
Cmca No.64
Density (g/cm3)
1.828
1.88
2.36
2.69
2.1
1.5
0.34
1.8244
2.6
2.4
Bandgap (eV) Refractive index
Isotopes Although twenty-three isotopes of phosphorus are known[19] (all possibilities from 24P up to 46 P), only 31P, with spin 1/2, is stable and is therefore present at 100% abundance. The half-integer spin and high abundance of 31P make it useful for nuclear magnetic resonance studies of biomolecules, particularly DNA. Two radioactive isotopes of phosphorus have half-lives which make them useful for scientific experiments. 32P has a half-life of 14.262 days and 33P has a half-life of 25.34 days. Biomolecules can be "tagged" with a radioisotope to allow for the study of very dilute samples. Radioactive isotopes of phosphorus include •
32
P; a beta-emitter (1.71 MeV) with a half-life of 14.3 days which is used routinely in life-science laboratories, primarily to produce radiolabeled DNA and RNA probes, e.g. for use in Northern blots or Southern blots. Because the high energy beta particles produced penetrate skin and corneas, and because any 32P ingested, inhaled, or absorbed is readily incorporated into bone and nucleic acids, Occupational Safety and Health Administration in the United States, and similar institutions in other developed countries require that a lab coat, disposable gloves, and safety glasses or goggles be worn when working with 32 P, and that working directly over an open container be avoided in order to protect the eyes. Monitoring personal, clothing, and surface contamination is also required. In
Phosphorus addition, due to the high energy of the beta particles, shielding this radiation with the normally used dense materials (e.g. lead), gives rise to secondary emission of X-rays via a process known as Bremsstrahlung, meaning braking radiation. Therefore shielding must be accomplished with low density materials, e.g. Plexiglas, Lucite, plastic, wood, or [20] water. • 33P; a beta-emitter (0.25 MeV) with a half-life of 25.4 days. It is used in life-science laboratories in applications in which lower energy beta emissions are advantageous such as DNA sequencing.
Chemical properties See also Phosphorus compounds • • • • • •
Hydrides: PH3,P2H4 Halides: PBr5, PBr3, PCl3, PI3 Oxides:P4O6, P4O10 Sulfides: P4S6, P4S10 Acids: H3PO2, H3PO4 Phosphates: (NH4)3PO4, Ca3(PO4)2, FePO4, Fe3(PO4)2, Na3PO4, Ca(H2PO4)2, KH2PO4
• Phosphides: Ca3P2, GaP, Zn3P2 Cu3P • Organophosphorus and organophosphates: Lawesson's reagent, Parathion, Sarin, Soman, Tabun, Triphenyl phosphine, VX nerve gas
Chemical bonding Because phosphorus is just below nitrogen in the periodic table, the two elements share many of their bonding characteristics. For instance, phosphine, PH3, is an analogue of ammonia, NH3. Phosphorus, like nitrogen, is trivalent in this molecule. The "trivalent" or simple 3-bond view is the pre-quantum mechanical Lewis structure, which although somewhat of a simplification from a quantum chemical point of view, illustrates some of the distinguishing chemistry of the element. In quantum chemical valence bond theory, the valence electrons are seen to be in mixtures of four s and p atomic orbitals, so-called hybrids. In this view, the three unpaired electrons in the three 3p orbitals combine with the two electrons in the 3s orbital to form three electron pairs of opposite spin, available for the formation of three bonds. The remaining hybrid orbital contains two paired non-bonding electrons, which show as a lone pair in the Lewis structure. The phosphorus cation is very similar to the nitrogen cation. In the same way that nitrogen forms the tetravalent ammonium ion, phosphorus can form the tetravalent phosphonium ion, and form salts such as phosphonium iodide [PH4]+[I−]. Like other elements in the third or lower rows of the periodic table, phosphorus atoms can expand their valence to make penta- and hexavalent compounds. The phosphorus chloride molecule is an example. When the phosphorus ligands are not identical, the more electronegative ligands are located in the apical positions and the least electronegative ligands are located in the axial positions. With strongly electronegative ions, in particular fluorine, hexavalency as in PF6− occurs as well. This octahedral ion is isoelectronic with SF6. In the bonding the six octahedral sp3d2 hybrid atomic orbitals play an important role.
6
Phosphorus Before extensive computer calculations were feasible, it was generally assumed that the nearby d orbitals in the n = 3 shell were the obvious cause of the difference in binding between nitrogen and phosphorus (i.e., phosphorus had 3d orbitals available for 3s and 3p shell bonding electron hybridisation, but nitrogen did not). However, in the early eighties the German theoretical chemist Werner Kutzelnigg[21] found from an analysis of computer calculations that the difference in binding is more likely due to differences in character between the valence 2p and valence 3p orbitals of nitrogen and phosphorus, respectively. The 2s and 2p orbitals of first row atoms are localized in roughly the same region of space, while the 3p orbitals of phosphorus are much more extended in space. The violation of the octet rule observed in compounds of phosphorus is then due to the size of the phosphorus atom, and the corresponding reduction of steric hindrance between its ligands. In modern theoretical chemistry, Kutzelnigg's analysis is generally accepted. The simple Lewis structure for the trigonal bipyramidal PCl5 molecule contains five covalent bonds, implying a hypervalent molecule with ten valence electrons contrary to the octet rule. An alternate description of the bonding, however, respects the octet rule by using 3-centre-4-electron (3c-4e) bonds. In this model the octet on the P atom corresponds to six electrons which form three Lewis (2c-2e) bonds to the three equatorial Cl atoms, plus the two electrons in the 3-centre Cl-P-Cl bonding molecular orbital for the two axial Cl electrons. The two electrons in the corresponding nonbonding molecular orbital are not included because this orbital is localized on the two Cl atoms and does not contribute to the electron density on the phosphorus atom. (However, it should always be remembered that the octet rule is not some universal rule of chemical bonding, and while many compounds obey it, there are many elements to which it does not apply).
Phosphine, diphosphine and phosphonium salts Phosphine (PH3) and arsine (AsH3) are structural analogs with ammonia (NH3) and form pyramidal structures with the phosphorus or arsenic atom in the center bound to three hydrogen atoms and one lone electron pair. Both are colourless, ill-smelling, toxic compounds. Phosphine is produced in a manner similar to the production of ammonia. Hydrolysis of calcium phosphide, Ca3P2, or calcium nitride, Ca3N2 produces phosphine or ammonia, respectively. Unlike ammonia, phosphine is unstable and it reacts instantly with air giving off phosphoric acid clouds. Arsine is even less stable. Although phosphine is less basic than ammonia, it can form some phosphonium salts (like PH4I), analogs of ammonium salts, but these salts immediately decompose in water and do not yield phosphonium (PH4+) ions. Diphosphine (P2H4 or H2P-PH2) is an analog of hydrazine (N2H4) that is a colourless liquid which spontaneously ignites in air and can disproportionate into phosphine and complex hydrides.
Halides The trihalides PF3, PCl3, PBr3 and PI3 and the pentahalides, PCl5 and PBr5 are all known and mixed halides can also be formed. The trihalides can be formed simply by mixing the appropriate stoichiometric amounts of phosphorus and a halide. For safety reasons, however, PF3 is typically made by reacting PCl3 with AsF5 and fractional distillation because the direct reaction of phosphorus with fluorine can be explosive. The pentahalides, PX5, are synthesized by reacting excess halide with either elemental phosphorus or with the
7
Phosphorus
8
corresponding trihalide. Mixed phosphorus halides are unstable and decompose to form simple halides. Thus 5PF3Br2 decomposes into 3PF5 and 2PBr5.
Oxides and oxyacids Phosphorus(III) oxide, P4O6 (also called tetraphosphorus hexoxide) and phosphorus(IV) oxide, P4O10 (or tetraphosphorus decoxide) are acid anhydrides of phosphorus oxyacids and hence readily react with water. P4O10 is a particularly good dehydrating agent that can even remove water from nitric acid, HNO3. The structure of P4O6 is like that of P4 with an oxygen atom inserted between each of the P-P bonds. The structure of P4O10 is like that of P4O6 with the addition of one oxygen bond to each phosphorus atom via a double bond and protruding away from the tetrahedral structure. Phosphorous oxyacids can have acidic protons bound to oxygen atoms and nonacidic protons which are bonded directly to the phosphorus atom. Although many oxyacids of phosphorus are formed, only six are important (see table), and three of them, hypophosphorous acid, phosphorous acid and phosphoric acid are particularly important ones. Oxidation State
Formula
Name
Acidic Protons
Compounds
+1
H3PO2
hypophosphorous acid
1
acid, salts
+3
H3PO3
(ortho)phosphorous acid
2
acid, salts
+5
(HPO3)n
metaphosphoric acids
n
salts (n=3,4)
+5
H5P3O10
triphosphoric acid
3
salts
+5
H4P2O7
pyrophosphoric acid
4
acid, salts
+5
H3PO4
(ortho)phosphoric acid
3
acid, salts
History and discovery The name Phosphorus in Ancient Greece was the name for the planet Venus and is derived from the Greek words (φως = light, φορέω = carry) which roughly translates as light-bringer or light carrier.[13] (In Greek mythology, Hesperus (evening star) and Eosphorus (dawnbearer) are close homologues, and also associated with Phosphorus-the-planet). The first recorded production of elemental phosphorus was in 1674 or 1675 by the German alchemist Hennig Brand through a preparation of urine, which contains considerable quantities of dissolved phosphates from normal metabolism.[13] Working in Hamburg, Brand attempted to create the fabled Philosopher's stone through the distillation of some salts by evaporating urine, and in the process produced a white material that glowed in the dark and burned brilliantly. His process originally involved letting urine stand for days until it gave off a terrible smell. Then he boiled it down to a paste, heated this paste to a high temperature, and led the vapours through water where he hoped they would condense to gold. Instead, he obtained a white, waxy substance that glowed in the dark. Brand had discovered phosphorus, the first element discovered since antiquity. We now know that Brand produced ammonium sodium hydrogen phosphate, (NH4)NaHPO4. While the quantities were essentially correct (it took about 1,100 L of urine to make about 60 g of phosphorus), it was unnecessary to allow the urine to rot. Later scientists would discover that fresh urine yielded the same amount of phosphorus.
Phosphorus Since that time, phosphors and phosphorescence were used loosely to describe substances that shine in the dark without burning. However, as mentioned above, even though the term phosphorescence was originally coined as a term by analogy with the glow from oxidation of elemental phosphorus, is now reserved for another fundamentally different process—re-emission of light after illumination. Phosphorus was gradually recognized as a chemical element in its own right at the emergence of the atomic theory that gradually occurred in the late part of the 18th century and the early 19th century (see John Dalton for more history). Brand at first tried to keep the method secret,[22] but later sold the recipe for 200 thaler to D Krafft from Dresden,[13] who could now make it as well, and toured much of Europe with it, including England, where he met with Robert Boyle. The secret that it was made from urine leaked out and first Johann Kunckel (1630-1703) in Sweden (1678) and later Boyle in London (1680) also managed to make phosphorus. Boyle states that Krafft gave him no information as to the preparation of phosphorus other than that it was derived from "somewhat that belonged to the body of man". This gave Boyle a valuable clue, however, so that he, too, managed to make phosphorus and published the method of its manufacture.[13] Later he improved Brand's process by using sand in the reaction (still using urine as base material), 4 NaPO3 + 2 SiO2 + 10 C → 2 Na2SiO3 + 10 CO + P4 Robert Boyle was the first to use phosphorus to ignite sulfur-tipped wooden splints, forerunners of our modern matches, in 1680. In 1769 Johan Gottlieb Gahn and Carl Wilhelm Scheele showed that calcium phosphate (Ca3(PO4)2) is found in bones and they obtained phosphorus from bone ash. Antoine Lavoisier recognized phosphorus as an element in 1777. Bone ash was the major source of phosphorus until the 1840s. Phosphate rock, a mineral containing calcium phosphate, was first used in 1850 and following the introduction of the electric arc furnace in 1890 this became the only source of phosphorus. Phosphorus, phosphates and phosphoric acid are still obtained from phosphate rock. Phosphate rock is a major feedstock in the fertilizer industry. Early matches used white phosphorus in their composition, which was dangerous due to its toxicity. Murders, suicides and accidental poisonings resulted from its use. (An apocryphal tale tells of a woman attempting to murder her husband with white phosphorus in his food, which was detected by the stew giving off luminous steam).[7] In addition, exposure to the vapours gave match workers a necrosis of the bones of the jaw, the infamous "phossy jaw." When a safe process for manufacturing red phosphorus was discovered, with its far lower flammability and toxicity, laws were enacted, under the Berne Convention (1906), requiring its adoption as a safer alternative for match manufacture. [11]
Spelling and etymology According to the Oxford English Dictionary the correct spelling of the element is phosphorus. The word phosphorous is the adjectival form of the P3+ valence: so, just as sulfur forms sulfurous and sulfuric compounds, phosphorus forms phosphorous compounds (see, e.g., phosphorous acid) and P5+ valency phosphoric compounds (see, e.g., phosphoric acids and phosphates).
9
Phosphorus
Occurrence Due to its reactivity with air and many other oxygen-containing substances, phosphorus is not found free in nature but it is widely distributed in many different minerals. Phosphate rock, which is partially made of apatite (an impure tri-calcium phosphate mineral), is an important commercial source of this element. About 50 percent of the global phosphorus reserves are in the Arab nations.[23] Large deposits of apatite are located in China, Russia, Morocco, Florida, Idaho, Tennessee, Utah, and elsewhere. Albright and Wilson in the United Kingdom and their Niagara Falls plant, for instance, were using phosphate rock in the 1890s and 1900s from Connetable, Tennessee and Florida; by 1950 they were using phosphate rock mainly from Tennessee and North Africa.[12] In the early 1990s Albright and Wilson's purified wet phosphoric acid business was being adversely affected by phosphate rock sales by China and the entry of their long-standing Moroccan phosphate suppliers into the purified wet phosphoric acid business.[24] In 2007, at the current rate of consumption, the supply of phosphorus was estimated to run out in 345 years.[25] However, scientists are now claiming that a "Peak Phosphorus" will occur in 30 years and that "At current rates, reserves will be depleted in the next 50 to 100 years."[26]
Production White phosphorus was first made commercially, for the match industry in the 19th century, by distilling off phosphorus vapour from precipitated phosphates, mixed with ground coal or charcoal, which was heated in an iron pot, in retort.[27] The precipitated phosphates were made from ground-up bones that had been de-greased and treated with strong acids. Carbon monoxide and other flammable gases produced during the reduction process were burnt off in a flare stack. This process became obsolete in the late 1890s when the electric arc furnace was adapted to reduce phosphate rock.[28] [4] Calcium phosphate (phosphate rock), mostly mined in Florida and North Africa, can be heated to 1,200-1,500 °C with sand, which is mostly SiO2, and coke (impure carbon) to produce vaporized tetraphosphorus, P4, (mp. 44.2 C) which is subsequently condensed into a white powder under water to prevent oxidation. Even under water, white phosphorus is slowly converted to the more stable red phosphorus allotrope (mp. 597 C). Both the white and red allotropes of phosphorus are insoluble in water. The electric furnace method allowed production to increase to the point where phosphorus could be used in weapons of war.[7] [12] In World War I it was used in incendiaries, smoke screens and tracer bullets.[12] A special incendiary bullet was developed to shoot at hydrogen-filled Zeppelins over Britain (hydrogen being highly inflammable if it can be ignited).[12] During World War II, Molotov cocktails of benzene and phosphorus were distributed in Britain to specially selected civilians within the British resistance operation, for defence; and phosphorus incendiary bombs were used in war on a large scale. Burning phosphorus is difficult to extinguish and if it splashes onto human skin it has horrific effects (see precautions below).[11] Today phosphorus production is larger than ever. It is used as a precursor for various chemicals,[29] in particular the herbicide glyphosate sold under the brand name Roundup. Production of white phosphorus takes place at large facilities and it is transported heated in liquid form. Some major accidents have occurred during transportation, train derailments
10
Phosphorus
11
at Brownston, Nebraska and Miamisburg, Ohio led to large fires. The worst accident in recent times was an environmental one in 1968 when phosphorus spilled into the sea from a plant at Placentia Bay, Newfoundland.[30]
Applications
Match striking surface made of a mixture of red phosphorus, glue and ground glass. (The glass is used to increase the friction.) Widely used compounds
Use
Ca(H2PO4)2•H2O
Baking powder & fertilizers
CaHPO4•2H2O
Animal food additive, toothpowder
H3PO4
Manufacture of phosphate fertilizers
PCl3
Manufacture of POCl3 and pesticides
POCl3
Manufacturing plasticizer
P4S10
Manufacturing of additives and pesticides
Na5P3O10
Detergents
Phosphorus, being an essential plant nutrient, finds its major use as a constituent of fertilizers for agriculture and farm production in the form of concentrated phosphoric acids, which can consist of 70% to 75% P2O5. Global demand for fertilizers led to large increase in phosphate (PO43-) production in the second half of the 20th century. Due to the essential nature of phosphorus to living organisms, the low solubility of natural phosphorus-containing compounds, and the slow natural cycle of phosphorus, the agricultural industry is heavily reliant on fertilizers which contain phosphate, mostly in the form of superphosphate of lime. Superphosphate of lime is a mixture of two phosphate salts, calcium dihydrogen phosphate Ca(H2PO4)2 and calcium sulfate dihydrate CaSO4•2H2O produced by the reaction of sulfuric acid and water with calcium phosphate. • Phosphorus is widely used to make organophosphorus compounds, through the intermediates phosphorus chlorides and two phosphorus sulfides: phosphorus pentasulfide, and phosphorus sesquisulfide.[12] Organophosphorus compounds have many applications, including in plasticizers, flame retardants, pesticides, extraction agents, and water treatment. [11] • Phosphorus is also an important component in steel production, in the making of phosphor bronze, and in many other related products.
Phosphorus • Phosphates are utilized in the making of special glasses that are used for sodium lamps.[31] • Bone-ash, calcium phosphate, is used in the production of fine china. [31] • Sodium tripolyphosphate made from phosphoric acid is used in laundry detergents in [31] some countries, but banned for this use in others. • Phosphoric acid made from elemental phosphorus is used in food applications such as some soda beverages. The acid is also a starting point to make food grade phosphates.[12] These include mono-calcium phosphate which is employed in baking powder and sodium tripolyphosphate and other sodium phosphates[12] . Among other uses these are used to improve the characteristics of processed meat and cheese. Others are used in toothpaste.[12] Trisodium phosphate is used in cleaning agents to soften water and for preventing pipe/boiler tube corrosion. • White phosphorus, called "WP" (slang term "Willie Peter") is used in military applications as incendiary bombs, for smoke-screening as smoke pots and smoke bombs, and in tracer ammunition. It is also a part of an obsolete M34 White Phosphorus US hand grenade. This multipurpose grenade was mostly used for signaling, smoke screens and inflammation; it could also cause severe burns and had a psychological impact on the enemy.[32] [33] • Red phosphorus is essential for manufacturing matchbook strikers, flares,[12] safety matches, pharmaceutical grade and street methamphetamine, and is used in cap gun caps. • Phosphorus sesquisulfide is used in heads of strike-anywhere matches.[12] • In trace amounts, phosphorus is used as a dopant for n-type semiconductors. • 32P and 33P are used as radioactive tracers in biochemical laboratories (see Isotopes).
Biological role Phosphorus is a key element in all known forms of life. Inorganic phosphorus in the form of the phosphate PO43- plays a major role in biological molecules such as DNA and RNA where it forms part of the structural framework of these molecules. Living cells also use phosphate to transport cellular energy in the form of adenosine triphosphate (ATP). Nearly every cellular process that uses energy obtains it in the form of ATP. ATP is also important for phosphorylation, a key regulatory event in cells. Phospholipids are the main structural components of all cellular membranes. Calcium phosphate salts assist in stiffening bones. [11]
Every cell has a membrane that separates it from its surrounding environment. Biological membranes are made from a phospholipid matrix and proteins, typically in the form of a bilayer. Phospholipids are derived from glycerol, such that two of the glycerol hydroxyl (OH) protons have been replaced with fatty acids as an ester, and the third hydroxyl proton has been replaced with phosphate bonded to another alcohol. [11] An average adult human contains a little less than 1 kg of phosphorus, about 85% of which is present in bones and teeth in the form of apatite, and the remainder inside cells in soft tissues. A well-fed adult in the industrialized world consumes and excretes about 1-3 g of phosphorus per day, with consumption in the form of inorganic phosphate and phosphorus-containing biomolecules such as nucleic acids and phospholipids; and excretion almost exclusively in the form of urine phosphate ion. Only about 0.1% of body phosphate circulates in the blood, but this amount reflects the amount of phosphate available to soft
12
Phosphorus tissue cells.
13 [11]
In medicine, low phosphate syndromes are caused by malnutrition, by failure to absorb phosphate, and by metabolic syndromes which draw phosphate from the blood (such as re-feeding after malnutrition) or pass too much of it into the urine. All are characterized by hypophosphatemia (see article for medical details), which is a condition of low levels of soluble phosphate levels in the blood serum, and therefore inside cells. Symptoms of hypophosphatemia include muscle and neurological dysfunction, and disruption of muscle and blood cells due to lack of ATP. Too much phosphate can lead to diarrhea and calcification (hardening) of organs and soft tissue, and can interfere with the body's ability to use iron, calcium, magnesium, and zinc.[34] Phosphorus edaphology phosphorus phosphorus phosphorus blooms.
is an essential macromineral for plants, which is studied extensively in in order to understand plant uptake from soil systems. In ecological terms, is often a limiting factor in many environments; i.e. the availability of governs the rate of growth of many organisms. In ecosystems an excess of can be problematic, especially in aquatic systems, see eutrophication and algal
Precautions Organic compounds of phosphorus form a wide class of materials, some of which are extremely toxic. Fluorophosphate esters are among the most potent neurotoxins known. A wide range of organophosphorus compounds are used for their toxicity to certain organisms as pesticides (herbicides, insecticides, fungicides, etc.) and weaponised as nerve agents. Most inorganic phosphates are relatively nontoxic and essential nutrients. For environmentally adverse effects of phosphates see eutrophication and algal blooms. [11] The white phosphorus allotrope should be kept under water at all times as it presents a significant fire hazard due to its extreme reactivity with atmospheric oxygen, and it should only be manipulated with forceps since contact with skin can cause severe burns. Chronic white phosphorus poisoning leads to necrosis of the jaw called "phossy jaw". Ingestion of white phosphorus may cause a medical condition known as "Smoking Stool Syndrome". [35] When the white form is exposed to sunlight or when it is heated in its own vapour to 250 °C, it is transmuted to the red form, which does not chemoluminesce in air. The red allotrope does not spontaneously ignite in air and is not as dangerous as the white form. Nevertheless, it should be handled with care because it reverts to white phosphorus in some temperature ranges and it also emits highly toxic fumes that consist of phosphorus oxides when it is heated. [11]
Phosphorus
Upon exposure to elemental phosphorus, in the past it was suggested to wash the affected area with 2% copper sulfate solution to form harmless compounds that can be washed away. According to the recent US Navy's Treatment of Chemical Agent Casualties and Conventional Military Chemical Injuries: FM8-285: Part 2 Conventional Military Chemical Injuries, "Cupric (copper(II)) sulfate has been used by U.S. personnel in the past and is still being used by some nations. However, copper sulfate is toxic and its use will be discontinued. Copper sulfate may produce kidney and cerebral toxicity as well as intravascular hemolysis."[36]
14
Phosphorus explosion
The manual suggests instead "a bicarbonate solution to neutralize phosphoric acid, which will then allow removal of visible white phosphorus. Particles often can be located by their emission of smoke when air strikes them, or by their phosphorescence in the dark. In dark surroundings, fragments are seen as luminescent spots." Then, "Promptly debride the burn if the patient's condition will permit removal of bits of WP which might be absorbed later and possibly produce systemic poisoning. DO NOT apply oily-based ointments until it is certain that all WP has been removed. Following complete removal of the particles, treat the lesions as thermal burns."[37] As white phosphorus readily mixes with oils, any oily substances or ointments are not recommended until the area is thoroughly cleaned and all white phosphorus removed.
US DEA List I status Phosphorus can reduce elemental iodine to hydroiodic acid, which is a reagent effective for reducing ephedrine or pseudoephedrine to methamphetamine.[38] For this reason, two allotropes of elemental phosphorus—red phosphorus and white phosphorus—were designated by the United States Drug Enforcement Administration as List I precursor chemicals under 21 CFR 1310.02 effective on November 17, 2001.[39] As a result, in the United States, handlers of red phosphorus or white phosphorus are subject to stringent regulatory controls pursuant to the Controlled Substances Act in order to reduce diversion of these substances for use in clandestine production of controlled substances.[39] [40] [41]
See also • White phosphorus (weapon)
References Notes [1] webelements (http:/ / www. webelements. com/ webelements/ compounds/ text/ P/ H4P2-13445506. html) [2] B. D. Ellis and C. L. B. Macdonald* "Phosphorus(I) Iodide: A Versatile Metathesis Reagent for the Synthesis of Low Oxidation State Phosphorus Compounds" Inorg. Chem., 2006, 45 (17), pp 6864 (http:/ / pubs. acs. org/ doi/ abs/ 10. 1021/ ic060186o) [3] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [4] Parkes and Mellor, pp. 718-720.
Phosphorus [5] " Nobel Prize in Chemistry 1956 - Presentation Speech by Professor A. Ölander (committee member) (http:/ / nobelprize. org/ chemistry/ laureates/ 1956/ press. html)". . Retrieved 2009-05-05. [6] " Phosphorus Topics page, at Lateral Science (http:/ / www. lateralscience. co. uk/ phos/ index. html)". . Retrieved 2009-05-05. [7] Emsley, John (2000). The Shocking History of Phosphorus. London: Macmillan. ISBN 0-330-39005-8. [8] Ana M. García-Campaña, Willy R. G. Baeyens (2001). Chemiluminescence in analytical chemistry (http:/ / books. google. com/ books?id=-EPSISOfCxYC). CRC Press. pp. 2–12. ISBN 0824704649. . [9] A. Holleman, N. Wiberg (1985). "XV 2.1.3". Lehrbuch der Anorganischen Chemie. de Gruyter. [10] Simon, Arndt (1997). "On the Polymorphism of White Phosphorus". Chemische Berichte 130: 1235. doi: 10.1002/cber.19971300911 (http:/ / dx. doi. org/ 10. 1002/ cber. 19971300911). [11] Lewis R. Goldfrank, Neal Flomenbaum, Mary Ann Howland, Robert S. Hoffman, Neal A. Lewin, Lewis S. Nelson (2006). Goldfrank's toxicologic emergencies (http:/ / books. google. com/ books?id=cvJuLqBxGUcC& pg=PA1487). McGraw-Hill Professional. pp. 1486–1489. ISBN 0071437630. . [12] Threlfall, R.E. (1951). 100 years of Phosphorus Making: 1851 - 1951. Oldbury: Albright and Wilson Ltd. [13] Parkes and Mellor, p. 717. [14] Parkes and Mellor, pp. 721-722. [15] A. Brown, S. Runquist (1965). "Refinement of the crystal structure of black phosphorus". Acta Crystallogr. 19: 684. doi: 10.1107/S0365110X65004140 (http:/ / dx. doi. org/ 10. 1107/ S0365110X65004140). [16] Cartz, L.;Srinivasa, S.R.;Riedner, R.J.;Jorgensen, J.D.;Worlton, T.G. (1979). "Effect of pressure on bonding in black phosphorus". Journal of Chemical Physics 71: 1718–1721. doi: 10.1063/1.438523 (http:/ / dx. doi. org/ 10. 1063/ 1. 438523). [17] Stefan Lange, Peer Schmidt, and Tom Nilges (2007). "Au3SnP7@Black Phosphorus: An Easy Access to Black Phosphorus". Inorg. Chem. 46: 4028. doi: 10.1021/ic062192q (http:/ / dx. doi. org/ 10. 1021/ ic062192q). [18] Piro, N. A. (2006). "Triple-Bond Reactivity of Diphosphorus Molecules". Science 313 (5791): 1276. doi: 10.1126/science.1129630 (http:/ / dx. doi. org/ 10. 1126/ science. 1129630). PMID 16946068. [19] " The Berkeley Laboratory Isotopes Project (http:/ / ie. lbl. gov/ education/ parent/ P_iso. htm)". . Retrieved 2009-05-05. [20] " Occupational Safety & Environmental Health: Phsophorus-21 (http:/ / www. oseh. umich. edu/ TrainP32. pdf)". . Retrieved 2009-05-05. [21] W. Kutzelnigg (1984). " Chemical Bonding in Higher Main Group Elements (http:/ / web. uvic. ca/ ~chem421/ ACIE_1984_Kutzelnigg_review. pdf)". Angewandte Chemie Int. (English) Ed. 23: 272–295. doi: 10.1002/anie.198402721 (http:/ / dx. doi. org/ 10. 1002/ anie. 198402721). . [22] J. M. Stillman (1960). The Story of Alchemy and Early Chemistry. New York: Dover. pp. 418–419. [23] " Phosphate Rock: Statistics and Information (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ phosphate_rock/ )". USGS. . Retrieved 2009-06-06. [24] Podger (2002), pages 297–298. [25] "How Long Will it Last?". New Scientist 194 (2605): 38–39. May 26, 2007. ISSN 4079 0262 4079 (http:/ / worldcat. org/ issn/ 0262). [26] Leo Lewis (2008-06-23). " Scientists warn of lack of vital phosphorus as biofuels raise demand (http:/ / business. timesonline. co. uk/ tol/ business/ industry_sectors/ natural_resources/ article4193017. ece)". The Times. . [27] Threlfall (1951), Chapter V: The foundations:1844–56: The phosphorus retort. [28] Threlfall (1951), Chapter VII: The second generation:1880–1915: Part 1: The electric furnace. [29] Aall C. H. (1952). "The American Phosphorus Industry". Industrial & Engineering Chemistry 44 (7): 1520–1525. doi: 10.1021/ie50511a018 (http:/ / dx. doi. org/ 10. 1021/ ie50511a018). [30] " ERCO and Long Harbour (http:/ / www. heritage. nf. ca/ law/ erco. html)". Memorial University of Newfoundland and the C.R.B. Foundation. . Retrieved 2009-06-06. [31] C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304814. [32] " Obsolete hand grenades (http:/ / www. globalsecurity. org/ military/ library/ policy/ army/ fm/ 3-23-30/ appe. htm)". GlobalSecurity.Org. . Retrieved 2009-08-03. [33] Dockery, Kevin (1997). Special Warfare Special Weapons. Chicago: Emperor's Press. ISBN 1-883-47600-3. [34] Anderson, John J. B. (01 Apr 1996). " Calcium, Phosphorus and Human Bone Development (http:/ / jn. nutrition. org/ cgi/ content/ abstract/ 126/ 4_Suppl/ 1153S)". Journal of Nutrition 126 (4 Suppl.): 1153S–1158S. PMID 8642449. . [35] " CBRNE - Incendiary Agents, White Phosphorus (Smoking Stool Syndrome) (http:/ / www. emedicine. com/ EMERG/ topic918. htm)". . Retrieved 2009-05-05. [36] " US Navy's Treatment of Chemical Agent Casualties and Conventional Military Chemical Injuries: FM8-285: Part 2 Conventional Military Chemical Injuries (http:/ / www. vnh. org/ FM8285/ Chapter/ chapter9. html)". .
15
Phosphorus Retrieved 2009-05-05. [37] This quote uses the word "phosphorescence", which is actually incorrect, WP, (White Phosphorous), exhibits chemoluminescence upon exposure to air and if there is any WP in the wound, covered by tissue or fluids such as blood serum, it will not chemoluminescence until it is moved to a position where the air can get at it and activate the chemoluminescent glow which requires a very dark room and dark adapted eyes to see clearly. [38] Skinner, H.F. (1990). "Methamphetamine synthesis via hydriodic acid/red phosphorus reduction of ephedrine". Forensic Science International 48 (2): 123–134. doi: 10.1016/0379-0738(90)90104-7 (http:/ / dx. doi. org/ 10. 1016/ 0379-0738(90)90104-7). [39] " 66 FR 52670—52675 (http:/ / frwebgate. access. gpo. gov/ cgi-bin/ getdoc. cgi?dbname=2001_register& docid=01-26013-filed)". 17 October 2001. . Retrieved 2009-05-05. [40] " 21 CFR 1309 (http:/ / www. access. gpo. gov/ nara/ cfr/ waisidx_06/ 21cfr1309_06. html)". . Retrieved 2009-05-05. [41] " 21 USC, Chapter 13 (Controlled Substances Act) (http:/ / www. usdoj. gov/ dea/ pubs/ csa. html)". . Retrieved 2009-05-05.
Sources • Emsley, John (2000). The Shocking history of Phosphorus. A biography of the Devil's Element. London: MacMillan. ISBN 0-333-76638-5. • Parkes, G.D. and Mellor, J.W. (1939). Mellor's Modern Inorganic Chemistry. London: Longman's Green and Co. • Podger, Hugh (2002). Albright & Wilson. The Last 50 years. Studley: Brewin Books. ISBN 1-85858-223-7. • Threlfall, Richard E. (1951). The Story of 100 years of Phosphorus Making: 1851–1951. Oldbury: Albright & Wilson ltd.
External links • Further warnings of toxic effects and recommendations for treatment can be found in " Emergency War Surgery NATO Handbook: Part I: Types of Wounds and Injuries: Chapter III: Burn Injury: Chemical Burns And White Phosphorus injury (http:/ / www. brooksidepress. org/ Products/ OperationalMedicine/ DATA/ operationalmed/ Manuals/ NATOEWS/ ch03/ 03ChemicalBurns. html)". http:/ / www. brooksidepress. org/ Products/ OperationalMedicine/ DATA/ operationalmed/ Manuals/ NATOEWS/ ch03/ 03ChemicalBurns. html. Retrieved 2009-05-05. • WebElements.com: Phosphorus (http:/ / www. webelements. com/ webelements/ elements/ text/ P/ index. html) • Simon, Fa; Pickering, Lk (Mar 1976). " Acute yellow phosphorus poisoning. "Smoking stool syndrome". (http:/ / toxnet. nlm. nih. gov/ cgi-bin/ sis/ search/ r?dbs+ hsdb:@term+ @rn+ 7723-14-0)" (Free full text). JAMA : the journal of the American Medical Association 235 (13): 1343–4. ISSN 0098-7484 (http:/ / worldcat. org/ issn/ 0098-7484). PMID 946251. http:/ / toxnet. nlm. nih. gov/ cgi-bin/ sis/ search/ r?dbs+ hsdb:@term+ @rn+ 7723-14-0. • eMedicine.com: Article on White Phophorus as used as weapon (http:/ / www. emedicine. com/ EMERG/ topic918. htm) • Website of the Technische Universität Darmstadt and the CEEP about Phosphorus Recovery (http:/ / www. phosphorus-recovery. tu-darmstadt. de)
16
Article Sources and Contributors
Article Sources and Contributors Phosphorus Source: http://en.wikipedia.org/w/index.php?oldid=307828753 Contributors: 131.111.8.xxx, 207.251.220.xxx, 21655, 2D, 88shrey, ABF, Achandrasekaran99, Achim1999, Adashiel, Addshore, Aeons, Ahkmemn, Ahoerstemeier, Aitias, Ajaxkroon, Alai, Alansohn, Albmont, Albo23, Anlace, Antandrus, Anwar saadat, AquaStreak, Arakunem, Archanamiya, Archimerged, Ardric47, Arnon Chaffin, Asmodeus Samael, Astrochemist, AuburnPilot, Axiosaurus, Axl, Axlenz, Ayla, BD2412, Bagel7, Bassbonerocks, Bbatsell, Bcorr, Beardo, Beetstra, BenM, Benbest, Bender235, Bensaccount, Benscripps, BillFlis, Biorga, Bkell, Black Walnut, Blastwizard, BlueEarth, Boboman828, Bongwarrior, Brian Crawford, Brian0918, BrianGV, Brianga, Bryan Derksen, Bunchofgrapes, Bushellman, Bushytails, CBDunkerson, CYD, Cadmium, Caltas, Can't sleep, clown will eat me, CanadianLinuxUser, Carnildo, Carom, Catbar, Catgut, Ceranthor, Ched Davis, Chekaz, Chemist1, Chemistryforlife, Chickydan, ChongDae, Chris Dybala, ChrisGriswold, Chvsanchez, Citanuleht, Closenplay, Cometstyles, Computer boy, Conversion script, Coppertwig, Cremepuff222, Cutler, DStoykov, DVD R W, DariusMazeika, Darrien, Darry2385, Darth Krayt, Darth Panda, David Latapie, Dbachmann, Dbtfz, Delta G, Demonator94, Deon, DerHexer, Dferg, Dfrg.msc, Dirac66, Dlae, Doczilla, Doovie, Doulos Christos, Download, DragonflySixtyseven, Dreish, Dumbo1, Durin, Ed42 311, Edgar181, Edsanville, Egomaniac, El C, Eldin raigmore, Element16, Emperorbma, Emre D., Enigmasoldier, Epbr123, Equendil, Eric119, Erik Zachte, Excirial, Felix Wan, Femto, Fett0001, Fieldday-sunday, FlavrSavr, Fluzwup, FlyingToaster, Fonzy, FreplySpang, Gaius Cornelius, Gegnome, Gene Nygaard, Geoseh, Giftlite, Glenn, Globe Collector, Gman124, Gracenotes, GraemeL, Grendelkhan, Gritchka, Gzuckier, Hak-kâ-ngìn, HappyCamper, Hdt83, Head, Helge Skjeveland, Hellbus, Hello Control, Heron, HiDrNick, Hieronymus Illinensis, Hu, IRD31416, Ian Pitchford, Ianweller, Icairns, Icseaturtles, Idamlaj, Imaninjapirate, ImperatorExercitus, InvictaHOG, Iridescent, Itub, J.delanoy, JForget, JFreeman, JSpung, Jaraalbe, Jazzman831, Jeffrey Mall, JimPAX, JimVC3, Jj137, Joebobway, John, John254, JohnyDog, Jons63, Jose77, JuniperBoy, Karl-Henner, Karlhahn, Keenan Pepper, Kevin Ryde, KnowledgeOfSelf, Kristen Eriksen, Kukini, Kurykh, Kwamikagami, Kx142, LarryMorseDCOhio, Leyo, LilHelpa, Looper5920, Lost Boy, LuigiManiac, Luna Santin, M1ss1ontomars2k4, M4gnum0n, MER-C, MPerel, MZMcBride, Mackeriv, Magus732, Mani1, Marek69, Materialscientist, Matnkat, Matticus78, Mausy5043, Mav, McBrainz, McNuggetsarecool, McSly, Mcpusc, Mentifisto, Michaeljay, Mifter, Mikaey, Mikeblas, Mikenorton, Mindspillage, Minesweeper, MissMJ, Montchav, Mooinglemur, Moomoomoo, Moonriddengirl, Moreschi, Mr Stephen, MrFish, Mrholybrain, Msps75, Mygerardromance, Nakon, Nathan, NawlinWiki, NeOak, Nergaal, Newone, Nickfield, Nickgetspwned, Night Gyr, Nihiltres, Nivix, Nonagonal Spider, Notchcode, Novangelis, NuclearWarfare, Nysin, Olin, Oneirist, Onevalefan, Optakeover, Oxymoron83, Ozzykhan, P. B. Mann, P.B. Pilhet, PRiis, Paleorthid, Patdoy3, Peachypoh, Persian Poet Gal, Petercorless, Peterlewis, Pgk, Pharaoh of the Wizards, Philip Trueman, Piano non troppo, Plasmic Physics, Ploober33, Plop, Pmish11, Poccil, Ponder, Poolkris, Poor Yorick, Psyche825, Pyrotec, Quadell, QueenCiti, Quintote, Qxz, RG2, RTC, RandomP, Realfoxxx, Redfarmer, Reguiieee, Remember, Res2216firestar, RexNL, Reyk, Riana, Richard L. Peterson, Richnotts, Rifleman 82, Rlove, RobertG, Rockstar915, Rogermw, Romanm, Rose Garden, Rror, Rumping, Rursus, RyanCross, Ryanaxp, SJP, Sanfranman59, Saperaud, Sbharris, Sceptre, Schneelocke, Sefog, Seinfreak37, Sengkang, Sfgagnon, Shaddack, Shadowin, Sikkema, Sintaku, Sir Nicholas de Mimsy-Porpington, Skatebiker, Sl, Smjg, Smokefoot, Smokizzy, SoWhy, Solipsist, Spellmaster, Spooser222, Squids and Chips, Ssri1983, StephP, Steven J. Anderson, Stifynsemons, Stone, Suisui, Sukrucetiner, Sunborn, SuperHamster, SweetNeo85, Synchronism, TempestSA, Tetracube, The Deviant, The Etceterist, The High Fin Sperm Whale, The Rambling Man, The Sanctuary Sparrow, Thehelpfulone, Thingg, Thricecube, Tim Starling, ToXiC, TomasBat, TreasuryTag, Trevor25, Treyd500, Trovatore, TwoOneTwo, Ufim, Uthbrian, V8rik, VASANTH S.N., Vampirehalfling, Vancouverguy, VasilievVV, Virek, Vladsinger, Vsmith, Vssun, WODUP, Waggers, Warut, Wereon, Wernher, Wikid77, Wimt, Windstreak7, Wtmitchell, Wyllium, XXI MR XRY IXx, Xaosflux, Xiankai, Xpyrda, Yachtsman1, Yamamoto Ichiro, Yyy, Zach4636, Zidane tribal, Zigger, Zomgadonggs, Ô, Александър, 1011 anonymous edits
Image Sources, Licenses and Contributors file:Unknown.svg Source: http://en.wikipedia.org/w/index.php?title=File:Unknown.svg License: Public Domain Contributors: Mav file:Electron shell 015 Phosphorus.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_015_Phosphorus.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 file:PhosphComby.jpg Source: http://en.wikipedia.org/w/index.php?title=File:PhosphComby.jpg License: GNU Free Documentation License Contributors: User:Maksim, User:Materialscientist Image:White phosphrous molecule.jpg Source: http://en.wikipedia.org/w/index.php?title=File:White_phosphrous_molecule.jpg License: Public Domain Contributors: Cadmium, Dzordzm, 1 anonymous edits File:Phosphorus-pentoxide-3D-balls.png Source: http://en.wikipedia.org/w/index.php?title=File:Phosphorus-pentoxide-3D-balls.png License: Public Domain Contributors: Benjah-bmm27 File:redPhosphorus.jpg Source: http://en.wikipedia.org/w/index.php?title=File:RedPhosphorus.jpg License: unknown Contributors: User:Materialscientist File:BlackPhosphorus.jpg Source: http://en.wikipedia.org/w/index.php?title=File:BlackPhosphorus.jpg License: unknown Contributors: User:Materialscientist Image:Match striking surface.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Match_striking_surface.jpg License: unknown Contributors: User:Startaq Image:Hazard F.svg Source: http://en.wikipedia.org/w/index.php?title=File:Hazard_F.svg License: Public Domain Contributors: BLueFiSH.as, MarianSigler, Matthias M., NielsF, Phrood, Pixeltoo, W!B:, 5 anonymous edits Image:Skull and crossbones.svg Source: http://en.wikipedia.org/w/index.php?title=File:Skull_and_crossbones.svg License: Public Domain Contributors: Andux, Bayo, Coyau, D0ktorz, Derbeth, Franzenshof, Ies, J.delanoy, Karelj, MarianSigler, Silsor, Stepshep, The Evil IP address, W!B:, 6 anonymous edits Image:Phosphorus explosion.gif Source: http://en.wikipedia.org/w/index.php?title=File:Phosphorus_explosion.gif License: Public Domain Contributors: US Gov.
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17
Sulfur
1
Sulfur phosphorus ← sulfur → chlorineO ↑ S ↓ Se
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
16S Periodic table
Appearance Lemon yellow crystals.
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Density (near r.t.) Density (near r.t.) Liquid density at m.p.Melting pointBoiling pointCritical pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
375
408
449
508
591
717
Sulfur
2
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 2252 kJ·mol−1 3rd: 3357 kJ·mol−1Covalent radiusVan der Waals radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityBulk modulusMohs hardnessCAS registry number Most stable isotopes Main article: Isotopes of sulfur iso
N.A.
half-life
32
95.02%
32
33
0.75%
33
34
4.21%
34
35
syn
87.32 d
36
0.02%
36
S S S S S
DM
DE (MeV)
DP
S is stable with 16 neutron S is stable with 17 neutron S is stable with 18 neutron β−
0.167
35
Cl
S is stable with 20 neutron
sulfur, S, 16 nonmetal16, 3, p32.065(5) g·mol−1 [Ne] 3s2 3p4 2, 8, 6 (Image) solid (alpha) 2.07 g·cm−3 (beta) 1.96 g·cm−3 (gamma) 1.92 g·cm−3 1.819 g·cm−3 388.36 K,115.21 °C,239.38 °F 717.8 K,444.6 °C,832.3 °F 1314 K, 20.7 MPa (mono) 1.727 kJ·mol−1 (mono) 45 kJ·mol−1 (25 °C) 22.75 J·mol−1·K−16, 5, 4, 3, 2, 1, -1, -2 (strongly acidic oxide) 2.58 (Pauling scale) 1st: 999.6 kJ·mol−1105±3 pm 180 pm orthorhombic diamagnetic[1] (20 °C) (amorphous) 2×1015Ω·m (300 K) (amorphous) 0.205 W·m−1·K−1 7.7 GPa 2.0 7704-34-9 Sulfur or sulphur (pronounced /ˈsʌlfər/, see spelling below) is the chemical element that has the atomic number 16. It is denoted with the symbol S. It is an abundant, multivalent non-metal. Sulfur in its native form is a yellow crystalline solid. In nature, it can be found as the pure element and as sulfide and sulfate minerals. It is an essential element for life and is found in two amino acids, cysteine and methionine. Its commercial uses are primarily in fertilizers, but it is also widely used in black gunpowder, matches, insecticides and fungicides. Elemental sulfur crystals are commonly sought after by mineral collectors for their brightly colored polyhedron shapes. In nonscientific contexts, it can also be referred to as brimstone.
History Sulfur (Sanskrit, sulvari; Latin sulfur or sulpur) was known in ancient times and is referred to in the Torah (Genesis).
Rough sulfur crystal
English translations of the Bible commonly referred to burning sulfur as "brimstone", giving rise to the name of 'fire-and-brimstone' sermons, in which listeners are reminded of the fate of eternal damnation that await the unbelieving and unrepentant. It is from this part of the Bible that Hell is implied to "smell of sulfur", although sulfur, in itself, is in fact odorless. The "smell
Sulfur
Sulfur crystal from Agrigento, Sicily.
3 of sulfur" usually refers to either the odor of hydrogen sulfide, e.g. from rotten egg, or of burning sulfur, which produces sulfur dioxide, the smell associated with burnt matches. The smell emanating from raw sulfur originates from a slow oxidation in the presence of air. Hydrogen sulfide is the principal odor of untreated sewage and is one of several unpleasant smelling sulfur-containing components of flatulence (along with sulfur-containing mercaptans).
A natural form of sulfur known as shiliuhuang was known in China since the 6th century BC and found in Hanzhong.[2] By the 3rd century, the Chinese discovered that sulfur could be extracted from pyrite.[2] Chinese Daoists were interested in sulfur's flammability and its reactivity with certain metals, yet its earliest practical uses were found in traditional Chinese medicine.[2] A Song Dynasty military treatise of 1044 AD described different formulas for Chinese black powder, which is a mixture of potassium nitrate (KNO3), charcoal, and sulfur. Early alchemists gave sulfur its own alchemical symbol which was a triangle at the top of a cross. In 1777, Antoine Lavoisier helped convince the scientific community that sulfur was an element and not a compound. In 1867, sulfur was discovered in underground deposits in Louisiana and Texas. The overlying layer of earth was quicksand, prohibiting ordinary mining operations; therefore, the Frasch process was developed.
Spelling and etymology The element has traditionally been spelled sulphur in the United Kingdom (since the 14th Century),[3] most of the Commonwealth including India, Malaysia, South Africa, and Hong Kong, along with the rest of the Caribbean and Ireland, but sulfur in the United States, while both spellings are used in Australia, New Zealand, Canada, and the Philippines. IUPAC adopted the spelling “sulfur” in 1990, as did the Royal Society of Chemistry Nomenclature Committee in 1992[4] and the Qualifications and Curriculum Authority for England and Wales recommended its use in 2000.[5] In Latin, the word is variously written sulpur, sulphur, and sulfur (the Oxford Latin Dictionary lists the spellings in this order). It is an original Latin name and not a Classical Greek loan, so the ph variant does not denote the Greek letter φ. Sulfur in Greek is thion (θείον), whence comes the prefix thio-. The simplification of the Latin words p or ph to an f appears to have taken place towards the end of the classical period, with the f spelling becoming dominant in the medieval period.[6] [7]
Sulfur
4
Characteristics At room temperature, sulfur is a soft, bright-yellow solid. Elemental sulfur has only a faint odor, similar to that of matches. The odor associated with rotten eggs is due to hydrogen sulfide (H2S) and organic sulfur compounds rather than elemental sulfur. Sulfur burns with a blue flame that emits sulfur dioxide, notable for its peculiar suffocating odor due to dissolving in the mucosa to form dilute sulfurous acid. Sulfur itself is insoluble in water, but soluble in carbon disulfide — and to a lesser extent in other non-polar organic solvents such as benzene and toluene. Common oxidation states of sulfur include −2, +2, +4 and +6. Sulfur forms stable compounds with all elements except the noble gases. Sulfur in the solid state ordinarily exists as cyclic crown-shaped S8 molecules.
Sulfur melts to a blood-red liquid. When burned, it emits a blue flame.
The crystallography of sulfur is complex. Depending on the specific conditions, the sulfur allotropes form several distinct crystal structures, with rhombic and monoclinic S8 best known. A noteworthy property of sulfur is that its viscosity in its molten state, unlike most other liquids, increases above temperatures of 200 °C due to the formation of polymers. The molten sulfur assumes a dark red color above this temperature. At higher temperatures, however, the viscosity is decreased as depolymerization occurs. Amorphous or "plastic" sulfur can be produced through the rapid cooling of molten sulfur. X-ray crystallography studies show that the amorphous form may have a helical structure with eight atoms per turn. This form is metastable at room temperature and gradually reverts back to crystalline form. This process happens within a matter of hours to days but can be rapidly catalyzed.
Allotropes
The structure of the cyclooctasulfur molecule, S8.
Sulfur forms more than 30 solid allotropes, more than any other element.[8] Besides S8, several other rings are known.[9] Removing one atom from the crown gives S7, which is more deeply yellow than S8. HPLC analysis of "elemental sulfur" reveals an equilibrium mixture of mainly S8, but also S7 and small amounts of S6.[10] Larger rings have been prepared, including S12 and S18.[11] [12] By contrast, sulfur's lighter neighbor oxygen only exists in two states of allotropic significance: O2 and O3. Selenium, the heavier analogue of sulfur, can form rings but is more often found as a polymer chain.
Sulfur
5
Isotopes Sulfur has 25 known isotopes, four of which are stable: 32S (95.02%), 33S (0.75%), 34S (4.21%), and 36S (0.02%). Other than 35S, the radioactive isotopes of sulfur are all short lived. 35S is formed from cosmic ray spallation of 40argon in the atmosphere. It has a half-life of 87 days. When sulfide minerals are precipitated, isotopic equilibration among solids and liquid may cause small differences in the δS-34 values of co-genetic minerals. The differences between minerals can be used to estimate the temperature of equilibration. The δC-13 and δS-34 of coexisting carbonates and sulfides can be used to determine the pH and oxygen fugacity of the ore-bearing fluid during ore formation. In most forest ecosystems, sulfate is derived mostly from the atmosphere; weathering of ore minerals and evaporites also contribute some sulfur. Sulfur with a distinctive isotopic composition has been used to identify pollution sources, and enriched sulfur has been added as a tracer in hydrologic studies. Differences in the natural abundances can also be used in systems where there is sufficient variation in the 34S of ecosystem components. Rocky Mountain lakes thought to be dominated by atmospheric sources of sulfate have been found to have different δS-34 values from lakes believed to be dominated by watershed sources of sulfate.
Occurrence Elemental sulfur can be found near hot springs and volcanic regions in many parts of the world, especially along the Pacific Ring of Fire. Such volcanic deposits are currently mined in Indonesia, Chile, and Japan. Sicily is also famous for its sulfur mines. Sulfur deposits are polycrystalline, and the largest documented single crystal measured 22x16x11 cm3.[13] [14] Significant deposits of elemental sulfur also exist in salt domes along the coast of the Gulf of Mexico, and in evaporites in eastern Europe and western Asia. The sulfur in these deposits is believed to come from the action of anaerobic bacteria on sulfate minerals, especially gypsum, although apparently native sulfur may be produced by geological processes alone, without the aid of living organisms (see below). However, fossil-based sulfur deposits from salt domes are the basis for commercial production in the United States, Poland, Russia, Turkmenistan, and Ukraine. Sulfur crystallites at Waiotapu hot springs, New Zealand
Sulfur
Sulfur production through hydrodesulfurization of oil, gas, and the Athabasca Oil Sands has produced a surplus — huge stockpiles of sulfur now exist throughout Alberta, Canada. Common naturally occurring sulfur compounds include the sulfide minerals, such as pyrite (iron sulfide), cinnabar (mercury sulfide), galena (lead sulfide), Sulfur recovered from hydrocarbons in sphalerite (zinc sulfide) and stibnite (antimony sulfide); Alberta, stockpiled for shipment at Vancouver, B.C. and the sulfates, such as gypsum (calcium sulfate), alunite (potassium aluminium sulfate), and barite (barium sulfate). It occurs naturally in volcanic emissions, such as from hydrothermal vents, and from bacterial action on decaying sulfur-containing organic matter. The distinctive colors of Jupiter's volcanic moon, Io, are from various forms of molten, solid and gaseous sulfur. There is also a dark area near the Lunar crater Aristarchus that may be a sulfur deposit. Sulfur is present in many types of meteorites. Ordinary chondrites contain on average 2.1% sulfur, and carbonaceous chondrites may contain as much as 6.6%. Sulfur in meteorites is normally present entirely as troilite (FeS), but other sulfides are found in some meteorites, and carbonaceous chondrites contain free sulfur, sulfates, and possibly other sulfur compounds.[15]
Extraction and production Extraction from natural resources Sulfur is extracted by mainly two processes: the Sicilian process and the Frasch process. The Sicilian process, which was first used in Sicily, was used in ancient times to get sulfur from rocks present in volcanic regions. In this process, the sulfur deposits are piled and stacked in brick kilns built on sloping hillsides, and with airspaces between them. Then powdered sulfur is put on top of the sulfur deposit and ignited. As the sulfur burns, the heat melts the sulfur deposits, causing the molten sulfur to flow down the sloping hillside. The molten sulfur can then be collected in wooden buckets. The second process used to obtain sulfur is the Frasch process. In this method, three concentric pipes are used: the outermost pipe contains superheated water, which melts the sulfur, and the innermost pipe is filled with hot compressed air, which serves to create foam and pressure. The resulting sulfur foam is then expelled through the middle pipe.[16] The Frasch process produces sulfur with a 99.5% purity content, and which needs no further purification. The sulfur produced by the Sicilian process must be purified by distillation.
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Production from hydrogen sulfide Chemically The Claus process is used to extract elemental sulfur from hydrogen sulfide produced in hydrodesulfurization of petroleum or from natural gas. Biologically In the biological route, hydrogen sulfide (H2S) from natural gas or refinery gas is absorbed with a slight alkaline solution in a wet scrubber. Or the sulfide is produced by biological sulfate reduction. In the subsequent process step, the dissolved sulfide is biologically converted to elemental sulfur. This solid sulfur is removed from the reactor. This process has been built on commercial scale. The main advantages of this process are: 1. 2. 3. 4. 5.
no use of expensive chemicals, the process is safe as the H2S is directly absorbed in an alkaline solution, no production of a polluted waste stream, re-usable sulfur is produced, and the process occurs under ambient conditions.
The biosulfur product is different from other processes in which sulfur is produced because the sulfur is hydrophillic. Next to straightforward reuses as source for sulfuric acid production, it can also be applied as sulfur fertilizer.[17]
Chemistry Inorganic compounds When dissolved in water, hydrogen sulfide is acidic and will react with metals to form a series of metal sulfides. Natural metal sulfides are common, especially those of iron. Iron sulfide is called pyrite, the so-called fool's gold. Pyrite can show semiconductor properties.[18] Galena, a naturally occurring lead sulfide, was the first semiconductor discovered, and found a use as a signal rectifier in the "cat's whiskers" of early crystal radios. Polymeric sulfur nitride has metallic properties even Sulfur powder. though it does not contain any metal atoms. This compound also has unusual electrical and optical properties. This polymer can be made from tetrasulfur tetranitride S4N4. Phosphorus sulfides are useful in synthesis. For example, P4S10 and its derivatives Lawesson's reagent and naphthalen-1,8-diyl 1,3,2,4-dithiadiphosphetane 2,4-disulfide are used to replace oxygen from some organic molecules with sulfur.
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• Sulfides (S2−), a complex family of compounds usually derived from S2−. Cadmium sulfide (CdS) is an example. • Sulfites (SO32−), the salts of sulfurous acid (H2SO3) which is generated by dissolving SO2 in water. Sulfurous acid and the corresponding sulfites are fairly strong reducing agents. Other compounds derived from SO2 include the pyrosulfite or metabisulfite ion (S2O52−).
• Sulfates (SO42−), the salts of sulfuric acid. Sulfuric acid also reacts with SO3 in equimolar ratios to form pyrosulfuric acid (H2S2O7). 2−
• Thiosulfates (S O 2
3
The sulfate anion, SO42−
). Sometimes referred as
thiosulfites or "hyposulfites", Thiosulfates are used in photographic fixing (HYPO) as reducing agents. Ammonium thiosulfate is being investigated as a cyanide replacement in leaching gold.[19] • Sodium dithionite, Na2S2O4, is the highly reducing dianion derived from hyposulfurous/dithionous acid. • Sodium dithionate (Na2S2O6). • Polythionic acids (H2SnO6), where n can range from 3 to 80. • Peroxymonosulfuric acid (H2SO5) and peroxydisulfuric acids (H2S2O8), made from the action of SO3 on concentrated H2O2, and H2SO4 on concentrated H2O2 respectively. • Sodium polysulfides (Na2Sx) • Sulfur hexafluoride, SF6, a dense gas at ambient conditions, is used as nonreactive and nontoxic propellant • Sulfur nitrides are chain and cyclic compounds containing only S and N. Tetrasulfur tetranitride S4N4 is an example. • Thiocyanates contain the SCN− group. Oxidation of thiocyanoate gives thiocyanogen, (SCN)2 with the connectivity NCS-SCN.
Organic compounds Many of the unpleasant odors of organic matter are based on sulfur-containing compounds such as methyl mercaptan and dimethyl sulfide. Thiols and sulfides are used in the odoriation of natural gas, notably, 2-methyl-2-propanethiol (t-butyl mercaptan). The odor of garlic and "skunk stink" are also caused by sulfur-containing organic compounds. Not all organic sulfur compounds smell unpleasant; for example, grapefruit mercaptan, a sulfur-containing monoterpenoid is responsible for the characteristic scent of grapefruit. It should be noted that this thiol is present in very low concentrations. In larger concentrations, the odor of this compound is that typical of all thiols, unpleasant. Sulfur-containing organic compounds include the following (R, R', and R are organic groups such as CH3):
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• Thioethers have the form R-S-R′. These compounds are the sulfur equivalents of ethers. • Sulfonium ions have the formula RR'S-'R'", i.e. where three groups are attached to the cationic sulfur center. Dimethylsulfoniopropionate (DMSP; (CH3)2S+CH2CH2COO−) is a sulfonium ion, which is important in the marine organic sulfur cycle. • Thiols (also known as mercaptans) have the form R-SH. These are the sulfur equivalents of alcohols. • Thiolates ions have the form R-S-. Such anions arise upon treatment of thiols with base. • Sulfoxides have the form R-S(=O)-R′. The simplest sulfoxide, DMSO, is a common solvent.
An organic sulfur compound, dithiane.
• Sulfones have the form R-S(=O)2-R′. A common sulfone is sulfolane C4H8SO2. See also Category: sulfur compounds and organosulfur chemistry
Applications One of the direct uses of sulfur is in vulcanization of rubber, where polysulfides crosslink organic polymers. Sulfur is a component of gunpowder. It reacts directly with methane to give carbon disulfide, which is used to manufacture cellophane and rayon.[20] Elemental sulfur is mainly used as a precursor to other chemicals. Approximately 85% (1989) is converted to sulfuric acid (H2SO4), which is of such prime importance to the world's economies that the production and consumption of sulfuric acid is an indicator of a nation's industrial development.[21] For example, more sulfuric acid is produced in the United States every year than any other industrial chemical. The principal use for the acid is the extraction of phosphate ores for the production of fertilizer manufacturing. Other applications of sulfuric acid include oil refining, wastewater processing, and mineral extraction.[20] Sulfur compounds are also used in detergents, fungicides, dyestuffs, and agrichemicals. In silver-based photography sodium and ammonium thiosulfate are used as "fixing agents." Sulfur is an ingredient in some acne treatments. An increasing application is as fertilizer. Standard sulfur is hydrophobic and therefore has to be covered with a surfactant by bacteria in the ground before it can be oxidized to sulfate. This makes it a slow release fertilizer, which cannot be taken up by the plants instantly, but has to be oxidized to sulfate over the growth season. Sulfur also improves the use efficiency of other essential plant nutrients, particularly nitrogen and phosphorus.[22] Biologically produced sulfur particles are naturally hydrophilic due to a biopolymer coating. This sulfur is therefore easier to disperse over the land (via spraying as a diluted slurry), and results in a faster release. Sulfites, derived from burning sulfur, are heavily used to bleach paper. They are also used as preservatives in dried fruit. Magnesium sulfate, better known as Epsom salts, can be used as a laxative, a bath additive, an exfoliant, a magnesium supplement for plants, or a desiccant.
Sulfur
Specialized applications Sulfur is used as a light-generating medium in the rare lighting fixtures known as sulfur lamps.
Historical applications In the late 18th century, furniture makers used molten sulfur to produce decorative inlays in their craft. Because of the sulfur dioxide produced during the process of melting sulfur, the craft of sulfur inlays was soon abandoned. Molten sulfur is sometimes still used for setting steel bolts into drilled concrete holes where high shock resistance is desired for floor-mounted equipment attachment points. Pure powdered sulfur was also used as a medicinal tonic and laxative. Sulfur was also used in baths for people who had fits.
Fungicide and pesticide Sulfur is one of the oldest fungicides and pesticides. Dusting sulfur, elemental sulfur in powdered form, is a common fungicide for grapes, strawberry, many vegetables and several other crops. It has a good efficacy against a wide range of powdery mildew diseases as well as black spot. In organic production, sulfur is the most important fungicide. It is the only fungicide used in organically farmed apple production against the main disease apple scab under colder conditions. Biosulfur (biologically produced elemental sulfur with hydrophillic characteristics) can be used well for these applications. Standard-formulation dusting sulfur is applied to crops with a sulfur duster or from a dusting plane. Wettable sulfur is the commercial name for dusting sulfur formulated with additional ingredients to make it water soluble. It has similar applications, and is used as a fungicide against mildew and other mold-related problems with plants and soil. Sulfur is also used as an "organic" (i.e. "green") insecticide (actually an acaricide) against ticks and mites. A common method of use is to dust clothing or limbs with sulfur powder. Some livestock owners set out a sulfur salt block as a salt lick.
Biological role See sulfur cycle for more on the inorganic and organic natural transformations of sulfur. Sulfur is an essential component of all living cells. Inorganic sulfur forms a part of iron-sulfur clusters, and sulfur is the bridging ligand in the CuA site of cytochrome c oxidase, a basic substance involved in utilization of oxygen by all aerobic life. Sulfur may also serve as chemical food source for some primitive organisms: some forms of bacteria use hydrogen sulfide (H2S) in the place of water as the electron donor in a primitive photosynthesis-like process in which oxygen is the electron receptor. The photosynthetic green and purple sulfur bacteria and some chemolithotrophs use elemental oxygen to carry out such oxidization of hydrogen sulfide to produce elemental sulfur (So), oxidation state = 0. Primitive bacteria which live around deep ocean volcanic vents oxidize hydrogen sulfide in this way with oxygen: see giant tube worm for an example of large organisms (via bacteria) making metabolic use of hydrogen sulfide as food to be oxidized. The so-called sulfur bacteria, by contrast, "breathe sulfate" instead of oxygen. They use sulfur as the electron acceptor, and reduce various oxidized sulfur compounds back into sulfide, often into hydrogen sulfide. They also can grow on a number of other partially
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Sulfur oxidized sulfur compounds (e. g. thiosulfates, thionates, polysulfides, sulfites). The hydrogen sulfide produced by these bacteria is responsible for the smell of some intestinal gases and decomposition products. Sulfur is a part of many bacterial defense molecules. For example, though sulfur is not a part of the lactam ring, it is a part of most beta lactam antibiotics, including the penicillins, cephalosporins, and monobactams. Sulfur is absorbed by plants via the roots from soil as the sulfate ion and reduced to sulfide before it is incorporated into cysteine and other organic sulfur compounds (see sulfur assimilation for details of this process). Sulfur is regarded as secondary nutrient although plant requirements for sulfur are equal to and sometimes exceed those for phosphorus. However sulfur is recognized as one of the major nutrients essential for plant growth, root nodule formation of legumes and plants protection mechanisms. Sulfur deficiency has become widespread in many countries in Europe. Because atmospheric inputs of sulfur will continue to decrease, the deficit in the sulfur input/output is likely to increase, unless sulfur fertilizers are used. In plants and animals the amino acids cysteine and methionine contain sulfur, as do all polypeptides, proteins, and enzymes which contain these amino acids. Homocysteine and taurine are other sulfur-containing acids which are similar in structure, but which are not coded for by DNA, and are not part of the primary structure of proteins. Glutathione is an important sulfur-containing tripeptide which plays a role in cells as a source of chemical reduction potential in the cell, through its sulfhydryl (-SH) moiety. Many important cellular enzymes use prosthetic groups ending with -SH moieties to handle reactions involving acyl-containing biochemicals: two common examples from basic metabolism are coenzyme A and alpha-lipoic acid. Disulfide bonds (S-S bonds) formed between cysteine residues in peptide chains are very important in protein assembly and structure. These strong covalent bonds between peptide chains give proteins a great deal of extra toughness and resiliency. For example, the high strength of feathers and hair is in part due to their high content of S-S bonds and their high content of cysteine and sulfur (eggs are high in sulfur because large amounts of the element are necessary for feather formation). The high disulfide content of hair and feathers contributes to their indigestibility, and also their odor when burned.
Traditional medical role for elemental sulfur In traditional medical skin treatment which predates modern era of scientific medicine, elemental sulfur has been used mainly as part of creams to alleviate various conditions such as psoriasis, eczema and acne. The mechanism of action is not known, although elemental sulfur does oxidize slowly to sulfurous acid, which in turn (though the action of sulfite) acts as a mild reducing and antibacterial agent.
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12
Precautions Carbon disulfide, carbon oxysulfide, hydrogen sulfide, and sulfur dioxide should all be handled with care. Although sulfur dioxide is sufficiently safe to be used as a food additive in small amounts, at high concentrations it reacts with moisture to form sulfurous acid which in sufficient quantities may harm the lungs, eyes or other tissues. In organisms without lungs such as insects or plants, it otherwise prevents respiration. Hydrogen sulfide is toxic. Although very pungent at first, it quickly deadens the sense of smell, so potential victims may be unaware of its presence until death or other symptoms occur.
Environmental impact The burning of coal and/or petroleum by industry and power plants generates sulfur dioxide (SO2), which reacts with atmospheric water and oxygen to produce sulfuric acid (H2SO4). This sulfuric acid is a component of acid rain, which lowers the pH of soil and freshwater bodies, sometimes resulting in substantial damage to the environment and chemical weathering of statues and structures. Fuel standards increasingly require sulfur to be extracted from fossil fuels to prevent the formation of acid rain. This extracted sulfur is then refined and represents a large portion of sulfur production. In coal fired power plants, the flue gases are sometimes purified. In more modern power plants that use syngas the sulfur is extracted before the gas is burned.
See also • • • • • • •
Sulfur cycle Stratospheric sulfur aerosols Disulfide bond Sulfonium S+, S+R3 Ultra-low sulfur diesel Claus process Shell-Paques sulfide removal/sulfur recovery process
References [1] Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf). CRC press. 2000. ISBN 0849304814. . [2] Zhang Yunming (1986). "The History of Science Society: Ancient Chinese Sulfur Manufacturing Processes". Isis 77: 487. doi: 10.1086/354207 (http:/ / dx. doi. org/ 10. 1086/ 354207). [3] http:/ / www. rod. beavon. clara. net/ sulphur. htm, retrieved 2nd April 2009 18:29 GMT. [4] Spelling of Sulfur (PDF) (http:/ / www. rsc. org/ delivery/ _ArticleLinking/ DisplayArticleForFree. cfm?doi=JM99101FP055& JournalCode=JM) [5] [6] [7] [8]
Worldwidewords (http:/ / www. worldwidewords. org/ topicalwords/ tw-sul1. htm), 9 December 2000. Vanderkrogt.net (http:/ / elements. vanderkrogt. net/ elem/ s. html). Kelly DP (1995) Sulfur and its Doppelgänger. Arch. Microbiol. 163: 157-158. Ralf Steudel, Bodo Eckert (2003). "Solid Sulfur Allotropes Sulfur Allotropes". Topics in Current Chemistry 230: 1–80. doi: 10.1007/b12110 (http:/ / dx. doi. org/ 10. 1007/ b12110).
[9] Steudel, R. (1982). "Homocyclic Sulfur Molecules". Topics Curr. Chem. 102: 149. [10] Tebbe, F. N.; Wasserman, E.; Peet, W. G.; Vatvars, A. and Hayman, A. C. (1982). "Composition of Elemental Sulfur in Solution: Equilibrium of S6, S7, and S8 at Ambient Temperatures". J. Am. Chem. Soc. 104: 4971. doi:
Sulfur 10.1021/ja00382a050 (http:/ / dx. doi. org/ 10. 1021/ ja00382a050). [11] Beat Meyer (1964). "Solid Allotropes of Sulfur". Chem. Rev. 64 (4): 429–451. doi: 10.1021/cr60230a004 (http:/ / dx. doi. org/ 10. 1021/ cr60230a004). [12] Beat Meyer (1976). "Elemental sulfur". Chem. Rev. 76: 367–388. doi: 10.1021/cr60301a003 (http:/ / dx. doi. org/ 10. 1021/ cr60301a003). [13] P. C. Rickwood (1981). " The largest crystals (http:/ / www. minsocam. org/ ammin/ AM66/ AM66_885. pdf)". American Mineralogist 66: 885-907. . [14] " The giant crystal project site (http:/ / giantcrystals. strahlen. org/ europe/ perticara. htm)". . Retrieved 2009-06-06. [15] B. Mason, Meteorites, (New York: John Wiley & Sons, 1962), p. 160. [16] Botsch, Walter (2001). "Chemiker, Techniker, Unternehmer: Zum 150. Geburtstag von Hermann Frasch" (in German). Chemie in unserer Zeit 35 (5): 324–331. doi: 10.1002/1521-3781(200110)35:5<324::AID-CIUZ324>3.0.CO;2-9 (http:/ / dx. doi. org/ 10. 1002/ 1521-3781(200110)35:5<324::AID-CIUZ324>3. 0. CO;2-9). [17] Zessen, E. van, et al. (2004). "Application of THIOPAQ(TM) biosulphur in agriculture". Proceedings of Sulphur 2004, Barcelona (Spain), 24 - 27 Oct. 57 - 68. [18] Nyle Steiner (22 February 1). " Iron Pyrites Negative Resistance Oscillator (http:/ / home. earthlink. net/ ~lenyr/ iposc. htm)". . Retrieved 2007-08-15. [19] http:/ / doccopper. tripod. com/ gold/ AltLixiv. html [20] Nehb, Wolfgang; Vydra, Karel (2006). "Sulfur". Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH Verlag. doi: 10.1002/14356007.a25_507.pub2 (http:/ / dx. doi. org/ 10. 1002/ 14356007. a25_507. pub2). [21] Sulfuric Acid Growth (http:/ / www. pafko. com/ history/ h_s_acid. html) [22] Sulfur as a fertilizer (http:/ / www. sulphurinstitute. org/ learnmore/ faq. cfm#plants)
Leslie KS, Millington GWM, Levell NJ. (2004) Sulphur and skin: from Satan to Saddam! J Cosm Dermatol 3: 94-98.
External links • Sulfur phase diagram (http:/ / library. tedankara. k12. tr/ chemistry/ vol2/ allotropy/ z129. htm) • WebElements.com – Sulfur (http:/ / www. webelements. com/ webelements/ elements/ text/ S/ index. html) • chemicalelements.com/sulfur (http:/ / www. chemicalelements. com/ elements/ s. html) • Crystalline, liquid and polymerization of sulphur on Vulcano Island, Italy (http:/ / www. stromboli. net/ perm/ vulcano/ sulphur-vulcano-en. html) • Sulfur and its use as a pesticide (http:/ / extoxnet. orst. edu/ pips/ sulfur. htm) • The Sulphur Institute (http:/ / www. sulphurinstitute. org/ )
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Article Sources and Contributors
Article Sources and Contributors Sulfur Source: http://en.wikipedia.org/w/index.php?oldid=307795402 Contributors: 203.109.250.xxx, 2D, A suyash, ABF, AWP1012933, Acroterion, Adamankin, Adamowen1, Adashiel, Adrian.benko, Ads85, Advanet, Afroman rmb, Ahoerstemeier, Aillema, Aitias, Alan Liefting, Alansohn, Alchie1, Ale jrb, Aleenf1, Alex.muller, AlexiusHoratius, AliaGemma, Allstarecho, AmiDaniel, Amzi, AncientToaster, Andre Engels, Andrew Maiman, Andrewjlockley, Andrij Kursetsky, Andycjp, AngelOfSadness, Angelic Wraith, Animum, Antonio Lopez, Archimerged, ArchonMagnus, Arsenal 14 8888, At the speed of light, Atlant, Aussie.mac95, AxelBoldt, Axlenz, BD2412, Backslash Forwardslash, Bassbonerocks, Bdesham, Bean159, Beetstra, Ben Arnold, Bender235, Benjah-bmm27, Benjiboi, Bergsten, Berkut, Bevanhouston, Bezking, Bhound89, Bit Lordy, BjKa, Blackfen, Bleedingshoes, Bloggeret, Bluap, BlueDevil, BlueEarth, Bluelip, Bobclay, Bobo192, Boccobrock, Bomac, Bongwarrior, Bonnocloudwolf, Bornhj, Boron1111, Branddobbe, Brandrewmiller, Briantw, Brockert, BrokenSphere, Brufydsy, Brutulf, Bryan Derksen, Buchanan-Hermit, Bwil, CWii, CYD, Cacycle, Calabraxthis, Caltas, Calvin 1998, CambridgeBayWeather, Can't sleep, clown will eat me, CanadianLinuxUser, CanisRufus, Capricorn42, CarbonCopy, Carnildo, Causesobad, Cazza619, Cbdorsett, Ccroberts, Cfailde, Chameleon, CharlotteWebb, Chasw0405, ChemNerd, Cherubfish, Chris 73, Chris Dybala, Cjh57, Clan Lord, Clondon, Closedmouth, Cometstyles, Confiteordeo, Conversion script, Coppertwig, Courtss, Crazy Boris with a red beard, CrazyChemGuy, Creamy Beaver99, Cremepuff222, Crosbiesmith, Crystal whacker, Cssiitcic, Cst17, Cureden, Curps, Cyrius, D13G054NCH3Z, DB, DINGGGGG!, DVD R W, DaBler, Daily Juice2, Damicatz, Dan D. 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15
Chlorine
1
Chlorine sulfur ← chlorine → argonF ↑ Cl ↓ Br
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
17Cl Periodic table
Appearance pale green gas
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensityMelting pointBoiling pointCritical pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
128
139
153
170
197
239
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 2298 kJ·mol−1 3rd: 3822 kJ·mol−1Covalent radiusVan der Waals radius
Chlorine
2
Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivitySpeed of soundCAS registry number Most stable isotopes Main article: Isotopes of chlorine iso
N.A.
half-life
35
75.77%
35
36
syn
3.01×105 y
Cl Cl
37
Cl
24.23%
DM
DE (MeV)
DP
Cl is stable with 18 neutron β−
0.709
36
ε
-
36
Ar S
37
Cl is stable with 20 neutron
chlorine, Cl, 17 Halogen17, 3, p35.453(2) g·mol−1 [Ne] 3s2 3p5 2, 8, 7 (Image) gas (0 °C, 101.325 kPa) 3.2 g/L 171.6 K,-101.5 °C,-150.7 °F 239.11 K,-34.4 °C,-29.27 °F 416.9 K, 7.991 MPa (Cl2) 6.406 kJ·mol−1 (Cl2) 20.41 kJ·mol−1 (25 °C) (Cl2) 33.949 J·mol−1·K−17, 6, 5, 4, 3, 2, 1, -1 (strongly acidic oxide) 3.16 (Pauling scale) 1st: 1251.2 kJ·mol−1102±4 pm 175 pm orthorhombic diamagnetic[1] (20 °C) > 10 Ω·m (300 K) 8.9x10-3 W·m−1·K−1 (gas, 0 °C) 206 m/s 7782-50-5 Chlorine (pronounced /ˈklɔərin/, from the Greek word 'χλωρóς' (khlôros, meaning 'pale green'), is the chemical element with atomic number 17 and symbol Cl. It is a halogen, found in the periodic table in group 17 (formerly VII, VIIa, or VIIb). As the chloride ion, which is part of common salt and other compounds, it is abundant in nature and necessary to most forms of life, including humans. In its elemental form (Cl2 or "dichlorine") under standard conditions, chlorine is a powerful oxidant and is used in bleaching and disinfectants. As a common disinfectant, chlorine compounds are used in swimming pools to keep them clean and sanitary. In the upper atmosphere, chlorine-containing molecules such as chlorofluorocarbons have been implicated in the destruction of the ozone layer.
Characteristics At standard temperature and pressure, two chlorine atoms form the diatomic molecule Cl2. This is a pale yellow-green gas that has its distinctive strong smell, the smell of bleach. The bonding between the two atoms is relatively weak (only of 242.580 ±0.004 kJ/mol) which makes the Cl2 molecule highly reactive. Along with fluorine, bromine, iodine and astatine, chlorine is a member of the halogen series that forms the group 17 of the periodic table—the most reactive group of elements. It combines readily with nearly all elements. Compounds with oxygen, nitrogen, xenon, and krypton are known, but do not form by direct reaction of the elements.[2] Chlorine, though very reactive, is not as extremely reactive as fluorine. Pure chlorine gas does, however, support combustion of organic compounds such as hydrocarbons, although the carbon component tends to burn incompletely, with much of it remaining as soot.[3] At 10 °C and atmospheric pressure, one liter of water dissolves 3.10 L of gaseous chlorine, and at 30 °C, 1 L of water dissolves only 1.77 liters of chlorine.[4] Chlorine is a member of the salt-forming halogen series and is extracted from chlorides through oxidation often by electrolysis. With metals, it forms salts called chlorides. As the
Chlorine
3
chloride ion, Cl−, it is also the most abundant dissolved ion in ocean water.
Isotopes Chlorine has isotopes with mass numbers ranging from 32 to 40. There are two principal stable isotopes, 35Cl (75.77%) and 37Cl (24.23%), giving chlorine atoms in bulk an apparent atomic weight of 35.4527 g/mol. Trace amounts of radioactive 36Cl exist in the environment, in a ratio of about 7x10−13 to 1 with stable isotopes. 36Cl is produced in the atmosphere by spallation of 36Ar by interactions with cosmic ray protons. In the subsurface environment, 36Cl is generated primarily as a result of neutron capture by 35Cl or muon capture by 40Ca. 36Cl decays to 36S and to 36Ar, with a combined half-life of 308,000 years. The half-life of this hydrophilic nonreactive isotope makes it suitable for geologic dating in the range of 60,000 to 1 million years. Additionally, large amounts of 36Cl were produced by irradiation of seawater during atmospheric detonations of nuclear weapons between 1952 and 1958. The residence time of 36Cl in the atmosphere is about 1 week. Thus, as an event marker of 1950s water in soil and ground water, 36Cl is also useful for dating waters less than 50 years before the present. 36Cl has seen use in other areas of the geological sciences, including dating ice and sediments.
Occurrence In nature, chlorine is found primarily as the chloride ion, a component of the salt that is deposited in the earth or dissolved in the oceans — about 1.9% of the mass of seawater is chloride ions. Even higher concentrations of chloride are found in the Dead Sea and in underground brine deposits. Most chloride salts are soluble in water, thus, chloride-containing minerals are usually only found in abundance in dry climates or deep underground. Common chloride minerals include halite (sodium chloride), sylvite (potassium chloride), and carnallite (potassium magnesium chloride hexahydrate). Over 2000 naturally-occurring organic chlorine compounds are known.[5] Industrially, elemental chlorine is usually produced by the electrolysis of sodium chloride dissolved in water. Along with chlorine, this chloralkali process yields hydrogen gas and sodium hydroxide, according to the following chemical equation: 2 NaCl + 2 H2O → Cl2 + H2 + 2 NaOH
Chlorine
4
History The most common compound of chlorine, sodium chloride, has been known since ancient times; archaeologists have found evidence that rock salt was used as early as 3000 BC and brine as early as 6000 BC.[6] The first compound of chlorine synthesized was probably hydrochloric acid (as a solution), which was prepared by the Persian alchemist Rhazes around 900 AD. Around 1200 AD, aqua regia (a mixture of nitric acid and hydrochloric acid) began to be used to dissolve gold, and today this is still one of the few reagents that will dissolve gold. Upon dissolving gold in aqua regia, chlorine gas is released along with other nauseating and irritating gases, but this wasn't known until much more recently. Chlorine was first prepared and studied in 1774 by Swedish chemist Carl Wilhelm Scheele, and therefore he is credited for its discovery.[7] He called it "dephlogisticated muriatic acid air" since it was a gas (then called "airs") and it came from hydrochloric acid (then known as "muriatic acid").[7] However, he failed to establish chlorine as an element, mistakenly thinking Liquid chlorine that it was the oxide obtained from the hydrochloric [7] acid (see phlogiston theory). He named the new element within this oxide as [7] muriaticum. Regardless of what he thought, Scheele did isolate chlorine by reacting MnO2 with HCl: 4 HCl + MnO2 → MnCl2 + 2 H2O + Cl2 Scheele observed several of the properties of chlorine: the bleaching effect on litmus, the deadly effect on insects, the yellow green color, and the smell similar to aqua regia. Claude Berthollet suggested that Scheele's dephlogisticated muratic acid air must be a combination of oxygen and an undiscovered element, muriaticum. In 1809 Joseph Louis Gay-Lussac and Louis-Jacques Thénard tried to decompose dephlogisticated muratic acid air by reacting it with charcoal to release the free element muriaticum (and carbon dioxide).[7] They did not succeed and published a report in which they considered the possibility that dephlogisticated muratic acid air is an element, but were not convinced.[8] In 1810, Sir Humphry Davy tried the same experiment again, and concluded that it was an element, and not a compound.[7] He named this new element as chlorine, from the Greek word χλωρος (chlōros), meaning green-yellow.[9] The name halogen, meaning salt producer, was originally defined for chlorine (in 1811 by Johann Salomo Christoph Schweigger), and it was later applied to the rest of the elements in this family. In 1822, Michael Faraday liquefied chlorine for the first time.[10] Chlorine was first used to bleach textiles in 1785.[11] In 1826, silver chloride was used to produce photographic images for the first time.[12] Chloroform was first used as an anesthetic in 1847.[12] Chlorine was first used as a germicide to prevent the spread of
Chlorine
5
puerperal fever in the maternity wards of Vienna General Hospital in Austria in 1847,[13] and in 1850 by John Snow to disinfect the water supply in London after an outbreak of cholera. The US Department of Treasury called for all water to be disinfected with chlorine by 1918.[12] Polyvinylchloride (PVC) was invented in 1912, initially without a purpose.[12] Chlorine gas was first introduced as a weapon on April 22, 1915 at Ypres by the German Army,[14] [15] and the results of this weapon were disastrous because gas masks had not yet been invented.
Production Gas extraction Chlorine can be manufactured by electrolysis of a sodium chloride solution (brine). The production of chlorine results in the co-products caustic soda (sodium hydroxide, NaOH) and hydrogen gas (H2). These two products, as well as chlorine itself, are highly reactive. Chlorine can also be produced by the electrolysis of a solution of potassium chloride, in which case the co-products are hydrogen and caustic potash (potassium hydroxide). There are three industrial methods for the extraction of chlorine by electrolysis of chloride solutions, all proceeding according to the following equations: Cathode: 2 H+ (aq) + 2 e− → H (g) 2
Anode: 2 Cl− (aq) → Cl (g) + 2 e− 2
Overall process: 2 NaCl (or KCl) + 2 H O → Cl + H 2
2
2
+ 2 NaOH (or KOH) Mercury cell electrolysis Mercury cell electrolysis, also known as the Castner-Kellner process, was the first method used at the end of the nineteenth century to produce chlorine on an industrial scale.[16] [17] The "rocking" Chlorine gas cells used have been improved over the years.[18] Today, in the "primary cell", titanium anodes (formerly graphite ones) are placed in a sodium (or potassium) chloride solution flowing over a liquid mercury cathode. When a potential difference is applied and current flows, chlorine is released at the titanium anode and sodium (or potassium) dissolves in the mercury cathode forming an amalgam. This flows continuously into a separate reactor ("denuder" or "secondary cell"), where it is usually converted back to mercury by reaction with water, producing hydrogen and sodium (or potassium) hydroxide at a commercially useful concentration (50% by weight). The mercury is then recycled to the primary cell. The mercury process is the least energy-efficient of the three main technologies (mercury, diaphragm and membrane) and there are also concerns about mercury emissions. It is estimated that there are still around 100 mercury-cell plants operating worldwide. In Japan, mercury-based chloralkali production was virtually phased out by 1987 (except for
Chlorine
6
the last two potassium chloride units shut down in 2003). In the United States, there will be only five mercury plants remaining in operation by the end of 2008. In Europe, mercury cells accounted for 43% of capacity in 2006 and Western European producers have committed to closing or converting all remaining chloralkali mercury plants by 2020.[19] Diaphragm cell electrolysis In diaphragm cell electrolysis, an asbestos (or polymer-fiber) diaphragm separates a cathode and an anode, preventing the chlorine forming at the anode from re-mixing with the sodium hydroxide and the hydrogen formed at the cathode.[20] This technology was also developed at the end of the nineteenth century. There are several variants of this process: the Le Sueur cell (1893), the Hargreaves-Bird cell (1901), the Gibbs cell (1908), and the Townsend cell (1904).[21] [22] The cells vary in construction and placement of the diaphragm, with some having the diaphragm in direct contact with the cathode. The salt solution (brine) is continuously fed to the anode compartment and flows through the diaphragm to the cathode compartment, where the caustic alkali is produced and the brine is partially depleted. As a result, diaphragm methods produce alkali that is quite dilute (about 12%) and of lower purity than do mercury cell methods. But diaphragm cells are not burdened with the problem of preventing mercury discharge into the environment. They also operate at a lower voltage, resulting in an energy savings over the mercury cell method,[22] but large amounts of steam are required if the caustic has to be evaporated to the commercial concentration of 50%. Membrane cell electrolysis Development of this technology began in the 1970s. The electrolysis cell is divided into two "rooms" by a cation permeable membrane acting as an ion exchanger. Saturated sodium (or potassium) chloride solution is passed through the anode compartment, leaving at a lower concentration.[23] Sodium (or potassium) hydroxide solution is circulated through the cathode compartment, exiting at a higher concentration. A portion of the concentrated sodium hydroxide solution leaving the cell is diverted as product, while the remainder is diluted with deionized water and passed through the electrolysis apparatus again. This method is more efficient than the diaphragm cell and produces very pure sodium (or potassium) hydroxide at about 32% concentration, but requires very pure brine. Other electrolytic processes Although a much lower production scale is involved, electrolytic diaphragm and membrane technologies are also used industrially to recover chlorine from hydrochloric acid solutions, producing hydrogen (but no caustic alkali) as a co-product. Furthermore, electrolysis of fused chloride salts (Downs process) also enables chlorine to be produced, in this case as a by-product of the manufacture of metallic sodium or magnesium.
Chlorine
7
Other methods Before electrolytic methods were used for chlorine production, the direct oxidation of hydrogen chloride with oxygen or air was exercised in the Deacon process: 4 HCl + O2 → 2 Cl2 + 2 H2O This reaction is accomplished with the use of copper(II) chloride (CuCl2) as a catalyst and is performed at high temperature (about 400 °C). The amount of extracted chlorine is approximately 80%. Due to the extremely corrosive reaction mixture, industrial use of this method is difficult and several pilot trials failed in the past. Nevertheless, recent developments are promising. Recently Sumitomo patented a catalyst for the Deacon process using ruthenium(IV) oxide (RuO2).[24] Another earlier process to produce chlorine was to heat brine with acid and manganese dioxide. 2 NaCl + 2 H2SO4 + MnO2 → Na2SO4 + MnSO4 + 2 H2O + Cl2 Using this process, chemist Carl Wilhelm Scheele was the first to isolate chlorine in a laboratory. The manganese can be recovered by the Weldon process.[25] In the latter half of the 19th century, prior to the adoption of electrolytic methods of chlorine production, there was substantial production of chlorine by these reactions to meet demand for bleach and bleaching powder for use by textile industries; by the 1880s the UK, as well as supporting its own (then not inconsiderable) domestic textile production was exporting 70,000 tons per year of bleaching powder.[26] This demand was met by capturing hydrochloric acid driven off as a gas during the production of alkali by the Leblanc process, oxidizing this to chlorine (originally by reaction with manganese dioxide), later by direct oxidation by air using the Deacon process (in which case impurities capable of poisoning the catalyst had first to be removed), and subsequently absorbing the chlorine onto lime. Small amounts of chlorine gas can be made in the laboratory by putting concentrated hydrochloric acid in a flask with a side arm and rubber tubing attached. Manganese dioxide is then added and the flask stoppered. The reaction is not greatly exothermic. As chlorine is denser than air, it can be collected by placing the tube inside a flask where it will displace the air. Once full, the collecting flask can be stoppered. Another method for producing small amounts of chlorine gas in a lab is by adding concentrated hydrochloric acid (typically about 5M) to sodium hypochlorite or sodium chlorate solution.
Chlorine
8
Industrial production Large-scale production of chlorine involves several steps and many pieces of equipment. The description below is typical of a membrane plant. The plant also simultaneously produces sodium hydroxide (caustic soda) and hydrogen gas. A typical plant consists of brine production/treatment, cell operations, chlorine cooling & drying, chlorine compression & liquefaction, liquid chlorine storage & loading, caustic handling, evaporation, storage & loading and hydrogen handling. Brine
Liquid Chlorine Analysis
Key to the production of chlorine is the operation of the brine saturation/treatment system. Maintaining a properly saturated solution with the correct purity is vital, especially for membrane cells. Many plants have a salt pile which is sprayed with recycled brine. Others have slurry tanks that are fed raw salt.
The raw brine is partially or totally treated with sodium hydroxide, sodium carbonate and a flocculant to reduce calcium, magnesium and other impurities. The brine proceeds to a large clarifier or a filter where the impurities are removed. The total brine is additionally filtered before entering ion exchangers to further remove impurities. At several points in this process, the brine is tested for hardness and strength. After the ion exchangers, the brine is considered pure, and is transferred to storage tanks to be pumped into the cell room. Brine, fed to the cell line, is heated to the correct temperature to control exit brine temperatures according to the electrical load. Brine exiting the cell room must be treated to remove residual chlorine and control pH levels before being returned to the saturation stage. This can be accomplished via dechlorination towers with acid and sodium bisulfite addition. Failure to remove chlorine can result in damage to the cells. Brine should be monitored for accumulation of both chlorate anions and sulfate anions, and either have a treatment system in place, or purging of the brine loop to maintain safe levels, since chlorate anions can diffuse through the membranes and contaminate the caustic, while sulfate anions can damage the anode surface coating. Cell room The building that houses many electrolytic cells is usually called a cell room or cell house, although some plants are built outdoors. This building contains support structures for the cells, connections for supplying electrical power to the cells and piping for the fluids. Monitoring and control of the temperatures of the feed caustic and brine is done to control exit temperatures. Also monitored are the voltages of each cell which vary with the electrical load on the cell room that is used to control the rate of production. Monitoring and control of the pressures in the chlorine and hydrogen headers is also done via pressure control valves. Direct current is supplied via a rectified power source. Plant load is controlled by varying the current to the cells. As the current is increased, flow rates for brine and caustic and deionized water are increased, while lowering the feed temperatures. Cooling and drying Chlorine gas exiting the cell line must be cooled and dried since the exit gas can be over 80°C and contains moisture that allows chlorine gas to be corrosive to iron piping. Cooling
Chlorine
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the gas allows for a large amount of moisture from the brine to condense out of the gas stream. This reduces both the cooling requirements and feed flow of sulfuric acid required in the drying towers. Cooling also improves the efficiency of both the compression and the liquefaction stage that follows. Chlorine exiting is ideally between 18°C and 25°C. After cooling the gas stream passes through a series of towers with counter flowing sulfuric acid. The sulfuric acid is fed into the final tower at 98% and the first tower typically has a strength between 66% and 76% depending on materials of construction. These towers progressively remove any remaining moisture from the chlorine gas. After exiting the drying towers the chlorine is filtered to remove any remaining sulfuric acid. Compression and liquefaction Several methods of compression may be used: liquid ring, reciprocating, or centrifugal. The chlorine gas is compressed at this stage and may be further cooled by inter- and after-coolers. After compression it flows to the liquefiers, where it is cooled enough to liquefy. Non condensable gases and remaining chlorine gas are vented off as part of the pressure control of the liquefaction systems. These gases are routed to a gas scrubber, producing sodium hypochlorite, or used in the production of hydrochloric acid (by combustion with hydrogen) or ethylene dichloride (by reaction with ethylene). Storage and loading Liquid chlorine is typically gravity-fed to storage tanks. It can be loaded into rail or road tankers via pumps or padded with compressed dry gas. Caustic handling, evaporation, storage and loading Caustic, fed to the cell room flows in a loop that is simultaneously bled off to storage with a part diluted with deionized water and returned to the cell line for strengthening within the cells. The caustic exiting the cell line must be monitored for strength, to maintain safe concentrations. Too strong or too weak a solution may damage the membranes. Membrane cells typically produce caustic in the range of 30% to 33% by weight. The feed caustic flow is heated at low electrical loads to control its exit temperature. Higher loads require the caustic to be cooled, to maintain correct exit temperatures. The caustic exiting to storage is pulled from a storage tank and may be diluted for sale to customers who require weak caustic or for use on site. Another stream may be pumped into a multiple effect evaporator set to produce commercial 50% caustic. Rail cars and tanker trucks are loaded at loading stations via pumps. Hydrogen handling Hydrogen produced may be vented unprocessed directly to the atmosphere or cooled, compressed and dried for use in other processes on site or sold to a customer via pipeline, cylinders or trucks. Some possible uses include the manufacture of hydrochloric acid or hydrogen peroxide, as well as desulfurization of petroleum oils, or use as a fuel in boilers or fuel cells. In Porsgrunn the byproduct is used for the hydrogen fueling station at Hynor. Energy consumption Production of chlorine is extremely energy intensive.[27] Energy consumption per unit weight of product is not far below that for iron and steel manufacture[28] and greater than for the production of glass[29] or cement.[30] Since electricity is an indispensable raw material for the production of chlorine, the energy consumption corresponding to the electrochemical reaction cannot be reduced. Energy savings arise primarily through applying more efficient technologies and reducing ancillary
Chlorine
10
energy use.
Compounds See also Chlorine compounds For general references to the chloride ion (Cl−), including references to specific chlorides, see chloride. For other chlorine compounds see chlorate (ClO−3), chlorite (ClO−2), hypochlorite (ClO−), and perchlorate (ClO−4), and chloramine (NH2Cl).[31] Other chlorine-containing compounds include: • Fluorides: chlorine monofluoride (ClF), chlorine trifluoride (ClF3), chlorine pentafluoride (ClF5) • Oxides: chlorine dioxide (ClO2), dichlorine monoxide (Cl2O), dichlorine heptoxide (Cl2O7) • Acids: hydrochloric acid (HCl), chloric acid (HClO3), and perchloric acid (HClO4)
Oxidation states Oxidation state
Name
Formula
Example compounds
−1
chlorides
Cl−
ionic chlorides, organic chlorides, hydrochloric acid
0
chlorine
Cl2
elemental chlorine
+1
hypochlorites
+3
chlorites
ClO−2
sodium chlorite
+5
chlorates
ClO−3
sodium chlorate, potassium chlorate, chloric acid
+7
perchlorates
ClO−4
potassium perchlorate, perchloric acid, magnesium perchlorate organic perchlorates, ammonium perchlorate
ClO−
sodium hypochlorite, calcium hypochlorite
Chlorine exists in all odd numbered oxidation states from −1 to +7, as well as the elemental state of zero. Progressing through the states, hydrochloric acid can be oxidized using manganese dioxide, or hydrogen chloride gas oxidized catalytically by air to form elemental chlorine gas. The solubility of chlorine in water is increased if the water contains dissolved alkali hydroxide. This is due to disproportionation: Cl2 + 2 OH− → Cl− + ClO− + H2O In hot concentrated alkali solution disproportionation continues: 2 ClO− → Cl− + ClO−2 ClO− + ClO−2 → Cl− + ClO−3 Sodium chlorate and potassium chlorate can be crystallized from solutions formed by the above reactions. If their crystals are heated, they undergo the final disproportionation step. 4 ClO−3 → Cl− + 3 ClO−4 This same progression from chloride to perchlorate can be accomplished by electrolysis. The anode reaction progression is:[32] Reaction
Electrode potential
Chlorine
11
Cl− + 2 OH− → ClO− + H2O + 2 e−
+0.89 volts
ClO− + 2 OH− → ClO−2 + H2O + 2 e−
+0.67 volts
ClO−2 + 2 OH− → ClO−3 + H2O + 2 e−
+0.33 volts
ClO−3 + 2 OH− → ClO−4 + H2O + 2 e−
+0.35 volts
Each step is accompanied at the cathode by 2 H2O + 2 e− → 2 OH− + H2 (−0.83 volts)
Applications and uses Production of industrial and consumer products Chlorine's principal applications are in the production of a wide range of industrial and consumer products.[33] [34] For example, it is used in making plastics, solvents for dry cleaning and metal degreasing, textiles, agrochemicals and pharmaceuticals, insecticides, dyestuffs, household cleaning products, etc.
Purification and disinfection Chlorine is an important chemical for water purification (such as water treatment plants), in disinfectants, and in bleach. Chlorine in water is more than three times more effective as a disinfectant against Escherichia coli than an equivalent concentration of bromine, and is more than six times more effective than an equivalent concentration of iodine.[35] Chlorine is usually used (in the form of hypochlorous acid) to kill bacteria and other microbes in drinking water supplies and public swimming pools. In most private swimming pools chlorine itself is not used, but rather sodium hypochlorite, formed from chlorine and sodium hydroxide, or solid tablets of chlorinated isocyanurates. Even small water supplies are now routinely chlorinated.[3] (See also chlorination) It is often impractical to store and use poisonous chlorine gas for water treatment, so alternative methods of adding chlorine are used. These include hypochlorite solutions, which gradually release chlorine into the water, and compounds like sodium dichloro-s-triazinetrione (dihydrate or anhydrous), sometimes referred to as "dichlor", and trichloro-s-triazinetrione, sometimes referred to as "trichlor". These compounds are stable while solid and may be used in powdered, granular, or tablet form. When added in small amounts to pool water or industrial water systems, the chlorine atoms hydrolyze from the rest of the molecule forming hypochlorous acid (HOCl) which acts as a general biocide killing germs, micro-organisms, algae, and so on.
Chlorine
Chemistry Elemental chlorine is an oxidizer. It undergoes halogen substitution reactions with lower halide salts. For example, chlorine gas bubbled through a solution of bromide or iodide anions oxidizes them to bromine and iodine respectively. Like the other halogens, chlorine participates in free-radical substitution reactions with hydrogen-containing organic compounds. This reaction is often—but not invariably—non-regioselective, and hence, may result in a mixture of isomeric products. It is often difficult to control the degree of substitution as well, so multiple substitutions are common. If the different reaction products are easily separated, e.g. by distillation, substitutive free-radical chlorination (in some cases accompanied by concurrent thermal dehydrochlorination) may be a useful synthetic route. Industrial examples of this are the production of methyl chloride, methylene chloride, chloroform and carbon tetrachloride from methane, allyl chloride from propylene, and trichloroethylene and tetrachloroethylene from 1,2-dichloroethane. Like the other halides, chlorine undergoes electrophilic additions reactions, most notably, the chlorination of alkenes and aromatic compounds with a Lewis acid catalyst. Organic chlorine compounds tend to be less reactive in nucleophilic substitution reactions than the corresponding bromine or iodine derivatives, but they tend to be cheaper. They may be activated for reaction by substituting with a tosylate group, or by the use of a catalytic amount of sodium iodide. Chlorine is used extensively in organic and inorganic chemistry as an oxidizing agent and in substitution reactions because chlorine often imparts many desired properties to an organic compound, due to its electronegativity. Chlorine compounds are used as intermediates in the production of a number of important commercial products that do not contain chlorine. Examples are: polycarbonates, polyurethanes, silicones, polytetrafluoroethylene, carboxymethyl cellulose and propylene oxide.
Use as a weapon • World War I Chlorine gas, also known as bertholite, was first used as a weapon in World War I by Germany on April 22, 1915 in the Second Battle of Ypres. As described by the soldiers it had a distinctive smell of a mixture between pepper and pineapple. It also tasted metallic and stung the back of the throat and chest. Chlorine can react with water in the mucosa of the lungs to form hydrochloric acid, an irritant which can be lethal. The damage done by chlorine gas can be prevented by a gas mask, or other filtration method, which makes the overall chance of death by chlorine gas much lower than those of other chemical weapons. It was pioneered by a German scientist later to be a Nobel laureate, Fritz Haber of the Kaiser Wilhelm Institute in Berlin, in collaboration with the German chemical conglomerate IG Farben, who developed methods for discharging chlorine gas against an entrenched enemy. It is alleged that Haber's role in the use of chlorine as a deadly weapon drove his wife, Clara Immerwahr, to suicide. After its first use, chlorine was utilized by both sides as a chemical weapon, but it was soon replaced by the more deadly gases phosgene and mustard gas.[36] • Iraq War
12
Chlorine Chlorine gas has also been used by insurgents against the local population and coalition forces in the Iraq War in the form of Chlorine bombs. On March 17, 2007, for example, three chlorine filled trucks were detonated in the Anbar province killing two and sickening over 350.[37] Other chlorine bomb attacks resulted in higher death tolls, with more than 30 deaths on two separate occasions.[38] Most of the deaths were caused by the force of the explosions rather than the effects of chlorine, since the toxic gas is readily dispersed and diluted in the atmosphere by the blast. The Iraqi authorities have tightened up security for chlorine, which is essential for providing safe drinking water for the population.
Chlorine cracking The element is widely used for purifying water owing to its powerful oxidizing properties, especially potable water supplies and water used in swimming pools. Several catastrophic collapses of swimming pool ceilings have occurred owing to stress corrosion cracking of stainless steel rods used to suspend them. Some polymers are also sensitive to attack, including acetal resin and polybutene. Both materials were used in hot and cold water domestic supplies, and stress Chlorine "attack" of an acetal resin plumbing joint. corrosion cracking caused widespread failures in the USA in the 1980s and '90s. One example shows an acetal joint in a water supply system, which when it fractured, caused substantial physical damage to computers in the labs below the supply. The cracks started at injection molding defects in the joint and grew slowly until finally triggered. The fracture surface shows iron and calcium salts which were deposited in the leaking joint from the water supply before failure.
Other uses Chlorine is used in the manufacture of numerous organic chlorine compounds, the most significant of which in terms of production volume are 1,2-dichloroethane and vinyl chloride, intermediates in the production of PVC. Other particularly important organochlorines are methyl chloride, methylene chloride, chloroform, vinylidene chloride, trichloroethylene, perchloroethylene, allyl chloride, epichlorohydrin, chlorobenzene, dichlorobenzenes and trichlorobenzenes. Chlorine is also used in the production of chlorates and in bromine extraction.
13
Chlorine
14
Health effects
NFPA 704
0 3 0 OX Chlorine is a toxic gas that irritates the respiratory system. Because it is heavier than air, it tends to accumulate at the bottom of poorly ventilated spaces. Chlorine gas is a strong oxidizer, which may react with flammable materials.[39] Chlorine is detectable in concentrations of as low as 1 ppm. Coughing and vomiting may occur at 30 ppm and lung damage at 60 ppm. About 1000 ppm can be fatal after a few deep breaths of the gas.[4] Breathing lower concentrations can aggravate the respiratory system, and exposure to the gas can irritate the eyes.[40] Chlorine's toxicity comes from its oxidizing power. When chlorine is inhaled at concentrations above 30ppm it begins to react with water and cells which change it into hydrochloric acid (HCl) and hypochlorous acid (HClO). When used at specified levels for water disinfection, although chlorine reaction with water itself usually doesn't represent a major concern for human health, other materials present in the water can generate disinfection by-products that can damage human health.[41] [42]
See also • Chloride • Polymer degradation
External links • Chlorine Institute [43] - Trade association and lobby group representing the interests of the chlorine industry • Chlorine Online [44] - Chlorine Online is an information resource produced by Eurochlor the business association of the European chlor-alkali industry • Electrolytic production [45] • Computational Chemistry Wiki [46] • Chlorine Production Using Mercury, Environmental Considerations and Alternatives [47] • National Pollutant Inventory - Chlorine [48] • National Institute for Occupational Safety and Health - Chlorine Page [49] • WebElements.com — Chlorine
[50]
Chlorine
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] Martha Windholz, editor ; Susan Budavari, associate editor ; Lorraine Y. Stroumtsos, assistant editor ; Margaret Noether Fertig, assistant editor. (1976). Merck Index of Chemicals and Drugs, 9th ed.. Rahway, N.J.: Merck & Co.. ISBN 0911910263. [3] C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304814. [4] " WebElements.com – Chlorine (http:/ / www. webelements. com/ webelements/ elements/ text/ Cl/ index. html)". Mark Winter [The University of Sheffield and WebElements Ltd, UK]. . Retrieved 2007-03-17. [5] " Risk assessment and the cycling of natural organochlorines (http:/ / www. eurochlor. org/ upload/ documents/ document236. pdf)". Euro Chlor. . Retrieved 2007-08-12. [6] " The earliest salt production in the world: an early Neolithic exploitation in Poiana Slatinei-Lunca, Romania (http:/ / antiquity. ac. uk/ ProjGall/ weller/ )". . Retrieved 2008-07-10. [7] " 17 Chlorine (http:/ / elements. vanderkrogt. net/ elem/ cl. html)". Elements.vanderkrogt.net. . Retrieved 2008-09-12. [8] Louis-Joseph Gay-Lussac, Louis-Jacques Thénard (1809). " On the nature and the properties of muriatic acid and of oxygenated muriatic acid (http:/ / web. lemoyne. edu/ ~giunta/ thenard. html)". Mémoires de Physique et de Chimie de la Société d'Arcueil 2: 339–358. . [9] Sir Humphry Davy (1811). " On a Combination of Oxymuriatic Gas and Oxygene Gas (http:/ / www. chemteam. info/ Chem-History/ Davy-Chlorine-1811. html)". Philosophical Transactions of the Royal Society 101: 155–162. doi: 10.1098/rstl.1811.0008 (http:/ / dx. doi. org/ 10. 1098/ rstl. 1811. 0008). . [10] " Discovery of Chlorine (http:/ / badley. info/ history/ Discovery-of-Chlorine-Great-Britain. event. html)". . Retrieved 2008-07-10. [11] " History of Chlorine (http:/ / members. aol. com/ manbio999/ chlorine. htm)". . Retrieved 2008-07-10. [12] Jacqueline Brazin. " Chlorine & its Consequences (http:/ / ocw. mit. edu/ NR/ rdonlyres/ Earth--Atmospheric--and-Planetary-Sciences/ 12-091January--IAP--2006/ 0EF9264B-3205-44A3-8306-8E8364917DF0/ 0/ brazin. pdf)". . Retrieved 2008-07-10. [13] " Chlorine Story (http:/ / www. americanchemistry. com/ s_chlorine/ sec_content. asp?CID=1166& DID=4476& CTYPEID=109)". americanchemistry. . Retrieved 2008-07-10. [14] " Chlorine - History (http:/ / www. drcordas. com/ education/ weaponsmassd/ Chlorine. pdf)". . Retrieved 2008-07-10. [15] " Weaponry: Use of Chlorine Gas Cylinders in World War I (http:/ / www. historynet. com/ weaponry-use-of-chlorine-gas-cylinders-in-world-war-i. htm)". historynet.com. . Retrieved 2008-07-10. [16] Pauling, Linus (1970). General Chemistry. Dover publications. ISBN 0-486-65622-5. [17] " Electrolytic Processes for Chlorine and Caustic Soda (http:/ / www. lenntech. com/ Chemistry/ electolytic-chlorine-caustic. htm)". Lenntech Water treatment & air purification Holding B.V., Rotterdamseweg 402 M, 2629 HH Delft, The Netherlands. . Retrieved 2007-03-17. [18] " Mercury cell (http:/ / www. eurochlor. org/ animations/ mercury-cell. asp)". Euro Chlor. . Retrieved 2007-08-15. [19] " Regional Awareness-raising Workshop on Mercury Pollution (http:/ / www. chem. unep. ch/ Mercury/ Pretoria-proceedings-finalPDFwriter. pdf)". UNEP. . Retrieved 2007-10-28. [20] " Diaphragm cell (http:/ / www. eurochlor. org/ animations/ diaphragm-cell. asp)". Euro Chlor. . Retrieved 2007-08-15. [21] " The Electrolysis of Brine (http:/ / www. saltsense. co. uk/ hist-chem12. htm)". Salt Manufacturers' Association. . Retrieved 2007-03-17. [22] Kiefer, David M.. " When the Industry Charged Ahead (http:/ / pubs. acs. org/ subscribe/ journals/ tcaw/ 11/ i04/ html/ 04chemistry. html)". Chemistry Chronicles. . Retrieved 2007-03-17. [23] " Membrane cell (http:/ / www. eurochlor. org/ animations/ membrane-cell. asp)". Euro Chlor. . Retrieved 2007-08-15. [24] Lopez, N (2008). "Mechanism of HCl oxidation (Deacon process) over RuO2". Journal of Catalysis 255: 29. doi: 10.1016/j.jcat.2008.01.020 (http:/ / dx. doi. org/ 10. 1016/ j. jcat. 2008. 01. 020). [25] " The Chlorine Industry (http:/ / www. lenntech. com/ Chemistry/ chlorine-industry. htm)". Lenntech Water treatment & air purification Holding B.V., Rotterdamseweg 402 M, 2629 HH Delft, The Netherlands. . Retrieved 2007-03-17. [26] Reader W J (1970 SBN 19 215937 2). Imperial Chemical Industries; A History. Volume 1. The Forerunners 1870-1926. Oxford University Press. p. 102. citing Haber L F (1958). The Chemical Industry during the Nineteenth Century. Oxford: Clarendon Press.
15
Chlorine [27] " Integrated Pollution Prevention and Control (IPPC) - Reference Document on Best Available Techniques in the Chlor-Alkali Manufacturing Industry (http:/ / www. jrc. es/ pub/ english. cgi/ d733217/ 05 Reference Document on Best Available Techniques in the Chlor-Alkali Manufacturing industry (adopted Dec 2001) - 5. 2 Mb)". European Commission. . Retrieved 2007-09-02. [28] " Integrated Pollution Prevention and Control (IPPC) - Best Available Techniques Reference Document on the Production of Iron and Steel (http:/ / www. jrc. es/ pub/ english. cgi/ d733208/ 02 Best Available Techniques Reference Document on the Production of Iron and Steel (adopted Dec 2001) - 9. 4Mb)". European Commission. . Retrieved 2007-09-02. [29] " Integrated Pollution Prevention and Control (IPPC) - Reference Document on Best Available Techniques in the Glass Manufacturing Industry (http:/ / www. jrc. es/ pub/ english. cgi/ d733226/ 08 Reference Document on Best Available Techniques in the Glass Manufacturing Industry (adopted Dec 2001) - 2. 7 Mb)". European Commission. . Retrieved 2007-09-02. [30] " Integrated Pollution Prevention and Control (IPPC) - Reference Document on Best Available Techniques in the Cement and Lime Manufacturing Industries (http:/ / www. jrc. es/ pub/ english. cgi/ d733211/ 03 Reference Document on Best Available Techniques in the Cement and Lime Manufacturing Industries (adopted Dec 2001) - 1. 3 Mb)". European Commission. . Retrieved 2007-09-02. [31] " Chlorine compounds of the month (http:/ / www. eurochlor. org/ index. asp?page=678)". Euro Chlor. . Retrieved 2007-08-29. [32] Cotton, F. Albert and Wilkinson, Geoffrey (1966). Advanced Inorganic Chemistry, 2nd ed.. John Wiley & sons. p. 568. [33] " Uses (http:/ / www. eurochlor. org/ uses)". Euro Chlor. . Retrieved 2007-08-20. [34] " Chlorine Tree (http:/ / www. chlorinetree. org)". Chlorine Tree. . Retrieved 2007-08-20. [35] Koski TA, Stuart LS, Ortenzio LF (1966). " Comparison of chlorine, bromine, iodine as disinfectants for swimming pool water (http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?tool=pubmed& pubmedid=4959984)". Applied Microbiology 14 (2): 276–279. PMID 4959984. . [36] " Weapons of War: Poison Gas (http:/ / www. firstworldwar. com/ weaponry/ gas. htm)". First World War.com. . Retrieved 2007-08-12. [37] Mahdi, Basim (2007-03-17). " Iraq gas attack makes hundreds ill (http:/ / www. cnn. com/ 2007/ WORLD/ meast/ 03/ 17/ iraq. main/ index. html)". CNN. . Retrieved 2007-03-17. [38] " 'Chlorine bomb' hits Iraq village (http:/ / news. bbc. co. uk/ 2/ hi/ middle_east/ 6660585. stm)". BBC News. 2007-05-17. . Retrieved 2007-05-17. [39] " Chlorine MSDS (http:/ / www. westlake. com/ datasheets/ MSDS_Chlorine. pdf)". October 23, 1997 (Revised November 1999. . [40] Chris Winder (2001). "The Toxicology of Chlorine". Environmental Research 85 (2): 105–114. doi: 10.1006/enrs.2000.4110 (http:/ / dx. doi. org/ 10. 1006/ enrs. 2000. 4110). [41] " What's in your Water?: Disinfectants Create Toxic By-products (http:/ / www. aces. uiuc. edu/ news/ stories/ news4724. html)". ACES News. College of Agricultural, Consumer and Environmental Sciences - University of Illinois at Urbana-Champaign. 2009-03-31. . Retrieved 2009-03-31. [42] Richardson, Sd; Plewa, Mj; Wagner, Ed; Schoeny, R; Demarini, Dm (Nov 2007). "Occurrence, genotoxicity, and carcinogenicity of regulated and emerging disinfection by-products in drinking water: a review and roadmap for research". Mutation research 636 (1-3): 178–242. doi: 10.1016/j.mrrev.2007.09.001 (http:/ / dx. doi. org/ 10. 1016/ j. mrrev. 2007. 09. 001). PMID 17980649. edit (http:/ / en. wikipedia. org/ wiki/ Template:cite_pmid/ 17980649) [43] [44] [45] [46] [47] [48] [49] [50]
http:/ / www. chlorineinstitute. org/ http:/ / www. eurochlor. org/ http:/ / electrochem. cwru. edu/ encycl/ art-b01-brine. htm http:/ / www. compchemwiki. org/ index. php?title=Cl2 http:/ / www. oceana. org/ chlorine http:/ / www. npi. gov. au/ database/ substance-info/ profiles/ 20. html http:/ / www. cdc. gov/ niosh/ topics/ chlorine/ http:/ / www. webelements. com/ webelements/ elements/ text/ Cl/ index. html
16
Article Sources and Contributors
Article Sources and Contributors Chlorine Source: http://en.wikipedia.org/w/index.php?oldid=307823011 Contributors: 000n, 0612, 12dstring, 213.106.152.xxx, 2L84MEANDU, A5b, ABF, APendleton, Aa35te, Ace of Spades IV, Acroterion, Adashiel, Adrian.benko, Ageofthewolf, Ahoerstemeier, Ajrocke, Alan Liefting, Alan Peakall, Alansohn, Ale jrb, Aljasm, Allstarecho, Alureiter, Alxndr, Anand Karia, Anastrophe, Andres, Andrewtindal, Anon0096, Antandrus, Anturiaethwr, Any Moose, Archimerged, Audirs8, Aussie Alchemist, AxelBoldt, Az1568, Bcorr, Beetstra, Bejnar, Benbest, Benjah-bmm27, Bennybp, Bhadani, Big Bird, Bighalonut, Biodragon, BlueEarth, BlueOrb, Bobblewik, Bomac, Bonadea, Bowlhover, Bradydavis, Brandonrush, Brian0918, Bryan Derksen, Bsadowski1, Bucketsofg, Bunchofgrapes, Buttonius, CL expert, CYD, Cacycle, Calabraxthis, Caleb303, Camw, Can't sleep, clown will eat me, CanisRufus, CardinalDan, Carnildo, Causesobad, Cgingold, Chad.netzer, CharlotteWebb, ChemGardener, Chocoforfriends, Chris Dybala, Chrislk02, Christophenstein, ClockworkSoul, Cmapm, Colbuckshot, Cometome11, Cometstyles, Computerjoe, Conversion script, Coolhandscot, Covington, DH85868993, DStoykov, Damicatz, Daniel,levine, DanielCD, Danny, Dantheman531, Darrien, Dave Runger, Daveh4h, Davehi1, David Latapie, Ddunn3233, DeadEyeArrow, Deano wig, Deor, DerHexer, Derek.cashman, Dina, Discospinster, Dogposter, Donarreiskoffer, Doonhamer, Dorftrottel, Doulos Christos, DrBob, Dureo, Dwmyers, Dycedarg, Dysepsion, Dzubint, ERcheck, ESkog, Edgar181, Edsanville, Efringe, El C, Eliashedberg, Elkman, Emhoo, Emperorbma, Emre D., Eric119, Erik Zachte, Espi, Evand, Everyking, Explicit, Farosdaughter, Fconaway, Femto, Fonzy, Footballfan190, Frankie0607, FranklinJ, Fraxxare, Fruge,
[email protected], Fvasconcellos, Gabethenerd, Gaius Cornelius, Galaxiaad, Gcsuchemistry, Gegnome, Genesis28, Giftlite, Gilliam, Gman124, GngstrMNKY, Gogo Dodo, GreatMizuti, Greenhorn1, Grendelkhan, Gustavb, Gwandoya, Gwernol, Gökhan, Hak-kâ-ngìn, Hall Monitor, Hankwang, Herbee, Heron, Hexagon1, Heyitsalexander, Hmains, Hobosteve, Homestarmy, INVERTED, IW.HG, Iamhim7789, Ian Pitchford, Icairns, Indiansprings, Iridescent, Isaac, Itfc+canes=me, Ithunn, Itub, IvoShandor, J Di, J-stan, J.delanoy, JForget, Jackollie, James086, Jaraalbe, Jayron32, Jeff G., Jimjamjak, Jirt, Jmundo, Joanjoc, Jodupouy, John, John Millikin, Jonathan tcn, Jordanp, Jordi.1991, Jose77, Jrugordon, Julesd, Justin00220, Kaelia, Kakofonous, Karlhahn, Karuna8, Kazikameuk, Keenan Pepper, Kelly Martin, Kenz0198, KevinCable, Kingcobra333, Kingpin13, Kirk Hilliard, Kntrabssi, KostasG, Kozuch, Krukowski, KuRiZu, Kukini, Kurykh, Kwamikagami, La Parka Your Car, LachlanA, Laughcosts, LeaveSleaves, LiDaobing, Lightmouse, LittleOldMe, Llort, Logical2u, Longhair, LorenzoB, LuigiManiac, Luk, Luna Santin, Lupinoid, Luxdormiens, Lychosis, MPerel, MSGJ, MSTCrow, MZMcBride, Magnus Manske, Majorly, Mani1, Manikwiki, Maozbj, Marc Venot, Materialscientist, Mathboy965, Mav, Maximus Rex, Mbessey, Mbeychok, McGoogalyBear, Meandmyself, Melaen, Mentifisto, Midgrid, Minesweeper, Mion, Mitchell1987, Mithgil, Mrdice, Mxn, Myproblem999, Naaa127, Nahat, Nakon, NawlinWiki, Neil916, Nergaal, Neverquick, NewEnglandYankee, Nibuod, Nihiltres, Nn123645, Noirum, NuclearWarfare, Nutriveg, Onco p53, Ortonmc, Oxymoron83, PFHLai, Paiev, PastaDruid, Patrick, Paul-L, Pb30, Pbijdens, Peter12220, PeterJeremy, Peterlewis, Peytonbland, Pgan002, Phgao, PhiJ, Philip Trueman, Piano non troppo, PierreAbbat, Pikasneez27, Pko, Poccil, Poolkris, Poor Yorick, Possum, Ppntori, Prissantenbär, Procellarum, Professor water, Prolog, Psyche825, Pusher, Quadell, Quinsareth, Quintote, Quistnix, Qxz, RAM, RTC, Razorflame, Rebel, Redsnork, Remember, Rettetast, RexNL, Riana, Richnotts, Rifleman 82, Rjccumbria, Rmhermen, Roberta F., Romanm, Rominandreu, RoyBoy, Ruddo2, Runnynose47, Rutherfordjigsaw, Ryan Postlethwaite, Ryan Roos, Ryanrulz 11, SEWilco, SJP, Sagaciousuk, Sairen42, Sam Korn, Sanfranman59, Saperaud, Scarian, Sceptre, Scetoaux, Schneelocke, Scott14, Sean William, SeanMack, Securiger, Seddon, Sengkang, Sfgagnon, Shaddack, Shimmin, Shoy, Shpoffo, Shrinkshooter, Shshshsh, Sikkema, Skatebiker, Sl, Slowking Man, Smokefoot, Sodiumjan, Soliloquial, Solipsist, Soobrickay, Spangineer, Spiffy sperry, Squids and Chips, Squirepants101, Starry maiden Gazer, Stephenb, Stone, Storm Rider, Subdolous, Subtractive, Suisui, Sunborn, Syrthiss, THEN WHO WAS PHONE?, Tadas12, Tarquin, Tarret, Tedickey, Teeeim, Teflex77, Tempodivalse, Tennekis, Tetracube, Texans37sox9, TheGrr, Thedreamdied, Themanoverthere, Thingg, Thinghy, Thomas H. Larsen, Thricecube, TigerShark, Timo25, Tiptoety, Tisdalepardi, Tj9991, Tom harrison, Tombomp, Tomtheman5, Tovkam, Tregoweth, Tresiden, Trevor MacInnis, Trevor223, Trigger820, Triona, Tristanb, Trougedoor122, Trusilver, Truthflux, Tttom, Tuspm, TutterMouse, Twas Now, Two hundred percent, Ugur Basak, Ultratomio, UnaLaguna, Until It Sleeps, Uyanga, VASANTH S.N., VX, Val42, Vancouverguy, Vandalwarrior, Versus22, Vsmith, Vssun, Waggers, Walkerma, Walkiped, Watch37264, WatermelonPotion, Wavelength, Wayward, Wiki alf, William Avery, Wimt, Wknight94, Wmahan, Workingclass91, Wpktsfs, Wrp103, Xerxes b, Yamamoto Ichiro, Yilloslime, Yyy, Zach4636, Zaui, 1192 anonymous edits
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17
Argon
1
Argon chlorine ← argon → potassiumNe ↑ Ar ↓ Kr
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
18Ar Periodic table
Appearance tasteless, odorless, colorless gas General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensityMelting pointBoiling pointTriple pointCritical pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
at T/K
10
100
1k
10 k
100 k
47
53
61
71
87
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 2665.8 kJ·mol−1 3rd: 3931 kJ·mol−1Covalent radiusVan der Waals radius Miscellaneous Crystal structureMagnetic orderingThermal conductivitySpeed of soundCAS registry number Most stable isotopes Main article: Isotopes of argon iso
N.A.
half-life
36
0.337%
36
37
syn
35 d
38
0.063%
38
39
syn
269 y
Ar Ar Ar Ar
DM
DE (MeV)
DP
Ar is stable with 18 neutron ε
0.813
37
β−
0.565
39
Cl
Ar is stable with 20 neutron K
Argon
2
40
99.600%
40
41
syn
109.34 min
β−
2.49
41
42
syn
32.9 y
β−
0.600
42
Ar Ar Ar
Ar is stable with 22 neutron K K
argon, Ar, 18 noble gases 18, 3, p39.948(1) g·mol−1 [Ne] 3s2 3p6 2, 8, 8 (Image) gas (0 °C, 101.325 kPa) 1.784 g/L 83.80 K,−189.35 °C,−308.83 °F 87.30 K,−185.85 °C,−302.53 °F 83.8058 K (-189°C), 69 kPa 150.87 K, 4.898 MPa 1.18 kJ·mol−1 6.43 kJ·mol−1 (25 °C) 20.786 J·mol−1·K−1 0 no data (Pauling scale) 1st: 1520.6 kJ·mol−1106±10 pm 188 pm face-centered cubic diamagnetic[1] (300 K) 17.72x10-3 W·m−1·K−1 (gas, 27 °C) 323 m/s 7440–37–1 Argon (pronounced /ˈɑrɡɒn/) is a chemical element designated by the symbol Ar. Argon has atomic number 18 and is the third element in group 18 of the periodic table (noble gases). Argon is present in the Earth's atmosphere at 0.94%. Terrestrially, it is the most abundant and most frequently used of the noble gases. Argon's full outer shell makes it stable and resistant to bonding with other elements. Its triple point temperature of 83.8058 K is a defining fixed point in the International Temperature Scale of 1990.
Cavendish's method for the isolation of argon. The gases are contained in a test-tube (A) standing over a large quantity of weak alkali (B), and the current is conveyed in wires insulated by U-shaped glass tubes (CC) passing through the liquid and round the mouth of the test-tube. The inner platinum ends (DD) of the wire receive a current from a battery of five Grove cells and a Ruhmkorff coil of medium size.
Argon
3
Characteristics Argon has approximately the same solubility in water as oxygen gas and is 2.5 times more soluble in water than nitrogen gas. Argon is colorless, odorless, tasteless and nontoxic in both its liquid and gaseous forms. Argon is inert under most conditions and forms no confirmed stable compounds at room temperature. Although argon is a noble gas, it has been found to have the capability of forming some compounds. For example, the creation of argon fluorohydride (HArF), a marginally stable compound of argon with A small piece of rapidly melting argon ice. fluorine and hydrogen, was reported by researchers at the University of Helsinki in 2000.[2] Although the neutral ground-state chemical compounds of argon are presently limited to HArF, argon can form clathrates with water when atoms of it are trapped in a lattice of the water molecules.[3] Also argon-containing ions and excited state complexes, such as ArH+ and ArF, respectively, are known to exist. Theoretical calculations have shown several argon compounds that should be stable but for which no synthesis routes are currently known.
History Argon (αργος, Greek meaning "inactive", in reference to its chemical inactivity)[4] [5] [6] was suspected to be present in air by Henry Cavendish in 1785 but was not isolated until 1894 by Lord Rayleigh and Sir William Ramsay in Scotland in an experiment in which they removed all of the oxygen, carbon dioxide, water and nitrogen from a sample of clean air.[7] [8] They had determined that nitrogen produced from chemical compounds was one-half percent lighter than nitrogen from the atmosphere. The difference seemed insignificant, but it was important enough to attract their attention for many months. They concluded that there was another gas in the air mixed in with the nitrogen.[9] Argon was also encountered in 1882 through independent research of H. F. Newall and W.N. Hartley. Each observed new lines in the color spectrum of air but were unable to identify the element responsible for the lines. Argon became the first member of the noble gases to be discovered. The symbol for argon is now Ar, but up until 1957 it was A.[10]
Occurrence Argon constitutes 0.934% by volume and 1.29% by mass of the Earth's atmosphere, and air is the primary raw material used by industry to produce purified argon products. Argon is isolated from air by fractionation, most commonly by cryogenic fractional distillation, a process that also produces purified nitrogen, oxygen, neon, krypton and xenon.[11] The Martian atmosphere in contrast contains 1.6% of argon-40 and 5 ppm of argon-36. The Mariner space probe fly-by of the planet Mercury in 1973 found that Mercury has a very thin atmosphere with 70% argon, believed to result from releases of the gas as a decay product from radioactive materials on the planet. In 2005, the Huygens probe also discovered the presence of argon-40 on Titan, the largest moon of Saturn.[12]
Argon
Isotopes The main isotopes of argon found on Earth are 40Ar (99.6%), 36Ar (0.34%), and 38Ar (0.06%). Naturally occurring 40K with a half-life of 1.25 × 109 years, decays to stable 40Ar (11.2%) by electron capture and positron emission, and also to stable 40Ca (88.8%) via beta decay. These properties and ratios are used to determine the age of rocks.[13] In the Earth's atmosphere, 39Ar is made by cosmic ray activity, primarily with 40Ar. In the subsurface environment, it is also produced through neutron capture by 39K or alpha emission by calcium. 37Ar is created from the decay of 40Ca as a result of subsurface nuclear explosions. It has a half-life of 35 days.[13]
Compounds Argon’s complete octet of electrons indicates full s and p subshells. This full outer energy level makes argon very stable and extremely resistant to bonding with other elements. Before 1962, argon and the other noble gases were considered to be chemically inert and unable to form compounds; however, compounds of the heavier noble gases have since been synthesized. In August 2000, the first argon compounds were formed by researchers at the University of Helsinki. By shining ultraviolet light onto frozen argon containing a small amount of hydrogen fluoride, argon fluorohydride (HArF) was formed.[2] [14] It is stable up to 40 kelvins (−233 °C).
Production Industrial Argon is produced industrially by the fractional distillation of liquid air, a process that separates liquid nitrogen, which boils at 77.3 K, from argon, which boils at 87.3 K and oxygen, which boils at 90.2 K. About 700,000 tons of argon are produced worldwide every year. [15] In radioactive decays 40
Ar, the most abundant isotope of argon, is produced by the decay of 40K with a half-life of 1.25 × 109 years by electron capture or positron emission. Because of this, it is used in potassium-argon dating to determine the age of rocks.
4
Argon
5
Applications There are several different reasons why argon is used in particular applications: • An inert gas is needed. In particular, argon is the cheapest alternative when diatomic nitrogen is not sufficiently inert. • Low thermal conductivity is required. • The electronic properties (ionization and/or the emission spectrum) are necessary.
Cylinders containing argon gas for use in extinguishing fire without damaging server equipment
Other noble gases would probably work as well in most of these applications, but argon is by far the cheapest. Argon is inexpensive since it is a byproduct of the production of liquid oxygen and liquid nitrogen, both of which are used on a large industrial scale. The other noble gases (except helium) are produced this way as well, but argon is the most plentiful since it has the highest concentration in the atmosphere. The bulk of argon applications arise simply because it is inert and relatively cheap.
Industrial processes Argon is used in some high-temperature industrial processes, where ordinarily non-reactive substances become reactive. For example, an argon atmosphere is used in graphite electric furnaces to prevent the graphite from burning. For some of these processes, the presence of nitrogen or oxygen gases might cause defects within the material. Argon is used in various types of metal inert gas welding such as tungsten inert gas welding, as well as in the processing of titanium and other reactive elements. An argon atmosphere is also used for growing crystals of silicon and germanium. Argon is an asphyxiant in the poultry industry, either for mass culling following disease outbreaks, or as a means of slaughter more humane than the electric bath. Argon's relatively high density causes it to remain close to the ground during gassing. Its non-reactive nature makes it suitable in a food product, and since it replaces oxygen within the dead bird, argon also enhances shelf life.[16] Argon is sometimes used for extinguishing fires where damage to equipment is to be avoided (see photo).
Argon
6
Preservative Argon is used to displace oxygenand moisture-containing air in packaging material to extend the shelf-lives of the contents. Aerial oxidation, hydrolysis, and other chemical reactions which degrade the products are retarded or prevented entirely. Bottles of high-purity chemicals and certain pharmaceutical products are available in sealed bottles or ampoules packed in argon. In wine making, argon is used to top-off barrels to avoid the aerial oxidation of ethanol to acetic acid during the aging process.
A sample of caesium is packed under argon to avoid reactions with air
Argon is also available in aerosol-type cans, which may be used to preserve compounds such as varnish, polyurethane, paint, etc. for storage after opening.[17] Since 2001 the American National Archives stores important national documents such as the Declaration of Independence and the Constitution within argon-filled cases to retard their degradation. Using argon reduces gas leakage, compared with the helium used in the preceding five decades. [18]
Laboratory equipment Argon may be used as the inert gas within Schlenk lines and gloveboxes. The use of argon over comparatively less expensive dinitrogen is preferred where nitrogen may react. Argon
may
be
used
as
the
carrier
gas
in
gas
chromatography and in electrospray ionization mass spectrometry; it is the gas of choice for the plasma used in ICP spectroscopy. Argon is preferred for the sputter coating of specimens for scanning electron microscopy. Argon ions are also used for sputtering in microelectronics.
Gloveboxes are typically filled with argon, which recirculate over scrubbers to maintain an oxygen- and moisture-free atmosphere
Medical use Cryosurgery procedures such as cryoablation use liquefied argon to destroy cancer cells. In surgery it is used in a procedure called "argon enhanced coagulation" which is a form of argon plasma beam electrosurgery. The procedure carries a risk of producing gas embolism in the patient and has resulted in the death of one person via this type of accident.[19] Blue argon lasers are used in surgery to weld arteries, destroy tumors, and to correct eye defects.[20] It has also used experimentally to replace nitrogen in the breathing or decompression mix, to speed the elimination of dissolved nitrogen from the blood.[21] See Argox (scuba).
Argon
7
Lighting Incandescent lights are filled with argon, to preserve the filaments at high temperature. It is used for the specific way it ionizes and emits light, such as in plasma globes and calorimetry in experimental particle physics. Gas-discharge lamps filled with argon provide blue light. Argon is also used for the creation of blue laser light.
Miscellaneous uses
An argon & mercury vapor discharge tube.
It is used for thermal insulation in energy efficient windows.[22] Argon is also used in technical scuba diving to inflate a dry suit, because it is inert and has low thermal conductivity.[23]
Compressed argon is allowed to expand, to cool the seeker heads of the AIM-9 Sidewinder missile, and other missiles that use cooled thermal seeker heads. The gas is stored at high pressure.[24] Argon-39, with a half-life of 269 years, has been used for a number of applications, primarily ice core and ground water dating. Also, potassium-argon dating is used in dating igneous rocks.
Safety Although argon is non-toxic, it does not satisfy the body's need for oxygen and is thus an asphyxiant. Argon is 25% more dense than air and is considered highly dangerous in closed areas. It is also difficult to detect because it is colorless, odorless, and tasteless. In confined spaces, it is known to result in death due to asphyxiation. A 1994 incident in Alaska that resulted in one fatality highlights the dangers of argon tank leakage in confined spaces, and emphasizes the need for proper use, storage and handling.[25]
Further reading • USGS Periodic Table - Argon
[26]
• Emsley, J., Nature’s Building Blocks; Oxford University Press: Oxford, NY, 2001; pp. 35–39. • Brown, T. L.; Bursten, B. E.; LeMay, H. E., In Chemistry: The Central Science, 10th ed.; Challice, J.; Draper, P.; Folchetti, N. et al.; Eds.; Pearson Education, Inc.: Upper Saddle River, NJ, 2006; pp. 276 and 289. • Triple point temperature: 83.8058 K - Preston-Thomas, H. (1990). "The International Temperature Scale of 1990 (ITS-90) [27]". Metrologia 27: 3–10. doi:10.1088/0026-1394/27/1/002 [28]. http:/ / www. bipm. org/ en/ publications/ its-90. html. • Triple point pressure: 69 kPa - "Section 4, Properties of the Elements and Inorganic Compounds; Melting, boiling, triple, and critical temperatures of the elements". CRC Handbook of Chemistry and Physics (85th ed.). Boca Raton, Florida: CRC Press. 2005.
Argon
8
External links • • • • •
WebElements.com – Argon [29] Diving applications: Why Argon? [30] Argon Ar Properties, Uses, Applications [31] Leftover Finish Preserver – Bloxygen [32] Periodic Table of the Elements: Argon [33]
pnb:نوگرآ
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] " HArF! Argon's not so noble after all - researchers make argon fluorohydride (http:/ / findarticles. com/ p/ articles/ mi_m1200/ is_9_158/ ai_65368548)". . [3] Belosludov, V. R.; O. S. Subbotin, D. S. Krupskii, O. V. Prokuda, and Y. Kawazoe (2006). "Microscopic model of clathrate compounds". J. Phys.: Conf. Ser. 29: 1. doi: 10.1088/1742-6596/29/1/001 (http:/ / dx. doi. org/ 10. 1088/ 1742-6596/ 29/ 1/ 001). [4] Hiebert, E. N. (1963). "In Noble-Gas Compounds". in Hyman, H. H.. Historical Remarks on the Discovery of Argon: The First Noble Gas. Chicago, IL: University of Chicago Press. pp. 3–20. [5] Travers, M. W. (1928). The Discovery of the Rare Gases. London: Edward Arnold & Co.. pp. 1–7. [6] Rayleigh, Lord; Ramsay, W. (1895). "Argon: A New Constituent of the Atmosphere". Chemical News 71,: 51–58. [7] Lord Rayleigh;William Ramsay (1894 - 1895). " Argon, a New Constituent of the Atmosphere (http:/ / www. jstor. org/ pss/ 115394)". Proceedings of the Royal Society of London 57 (1): 265–287. doi: 10.1098/rspl.1894.0149 (http:/ / dx. doi. org/ 10. 1098/ rspl. 1894. 0149). . [8] William Ramsay. " Nobel Lecture in Chemistry, 1904 (http:/ / nobelprize. org/ nobel_prizes/ chemistry/ laureates/ 1904/ ramsay-lecture. html)". . [9] " About Argon, the Inert; The New Element Supposedly Found in the Atmosphere (http:/ / query. nytimes. com/ gst/ abstract. html?res=9B04E3D61139E033A25750C0A9659C94649ED7CF)". The New York Times. . Retrieved 2009-02-01. [10] Holden, Norman E. (12). " History of the Origin of the Chemical Elements and Their Discoverers (http:/ / www. nndc. bnl. gov/ content/ elements. html)". National Nuclear Data Center (NNDC). . [11] " Argon, Ar (http:/ / elements. etacude. com/ Ar. php)". . Retrieved 2007-03-08. [12] " Seeing, touching and smelling the extraordinarily Earth-like world of Titan (http:/ / www. esa. int/ esaCP/ SEMHB881Y3E_index_0. html)". European Space Agency. 21. . [13] " 40Ar/39Ar dating and errors (http:/ / www. geoberg. de/ text/ geology/ 07011601. php)". . Retrieved 2007-03-07. [14] Bartlett, Neil. " The Noble Gases (http:/ / pubs. acs. org/ cen/ 80th/ noblegases. html)". Chemical & Engineering News. . [15] " Periodic Table of Elements: Argon – Ar (http:/ / environmentalchemistry. com/ yogi/ periodic/ Ar. html)". Environmentalchemistry.com. . Retrieved 2008-09-12. [16] D. L. Fletcher. "[Downbound.com Symposium: Recent Advances in Poultry Slaughter Technology Slaughter Technology]". Downbound.com. Retrieved 2009-08-01. [17] US Patent 6629402 [18] " Schedule for Renovation of the National Archives Building (http:/ / www. archives. gov/ press/ press-kits/ charters. html#pressrelaese1)". . Retrieved 2009-07-07. [19] " Fatal Gas Embolism Caused by Overpressurization during Laparoscopic Use of Argon Enhanced Coagulation (http:/ / www. mdsr. ecri. org/ summary/ detail. aspx?doc_id=8248)". MDSR. 24. . [20] Fujimoto, James; Rox Anderson, R. (2006). " Tissue Optics, Laser-Tissue Interaction, and Tissue Engineering (http:/ / www. spie. org/ Conferences/ Programs/ 06/ pw/ BiOSAbstracts. pdf)" (pdf). Biomedical Optics. pp. 77-88. . Retrieved 2007-03-08. [21] Pilmanis Andrew A, Balldin UI, Webb James T, Krause KM (December 2003). "Staged decompression to 3.5 psi using argon-oxygen and 100% oxygen breathing mixtures". Aviation, Space, Environmental Medicine 74 (12): 1243–50. PMID 14692466. [22] " Energy-Efficient Windows (http:/ / www. finehomebuilding. com/ how-to/ articles/ understanding-energy-efficient-windows. aspx)". FineHomebuilding.com. . Retrieved 2009-08-01.
Argon
9
[23] Nuckols ML, Giblo J, Wood-Putnam JL. (September 15-18, 2008). " Thermal Characteristics of Diving Garments When Using Argon as a Suit Inflation Gas. (http:/ / archive. rubicon-foundation. org/ 7962)". Proceedings of the Oceans 08 MTS/IEEE Quebec, Canada Meeting (MTS/IEEE). . Retrieved 2009-03-02. [24] " Description of Aim-9 Operation (http:/ / home. wanadoo. nl/ tcc/ rnlaf/ aim9. html)". planken.org. . Retrieved 2009-02-01. [25] Middaugh, John (1994-06-23). " Welder's Helper Asphyxiated in Argon-Inerted Pipe (FACE AK-94-012) (http:/ / www. hss. state. ak. us/ dph/ ipems/ occupation_injury/ reports/ docs/ 94ak012. htm)". State of Alaska Department of Public Health. . Retrieved 2009-02-01. [26] [27] [28] [29] [30] [31] [32] [33]
http:/ / wwwrcamnl. wr. usgs. gov/ isoig/ period/ ar_iig. html http:/ / www. bipm. org/ en/ publications/ its-90. html http:/ / dx. doi. org/ 10. 1088%2F0026-1394%2F27%2F1%2F002 http:/ / www. webelements. com/ webelements/ elements/ text/ Ar/ index. html http:/ / www. decompression. org/ maiken/ Why_Argon. htm http:/ / www. uigi. com/ argon. html http:/ / www. bloxygen. com http:/ / www. lenntech. com/ Periodic-chart-elements/ Ar-en. htm
Article Sources and Contributors
Article Sources and Contributors Argon Source: http://en.wikipedia.org/w/index.php?oldid=308216912 Contributors: 2D, ABF, Abce2, Abrech, Adashiel, Addshore, Ahoerstemeier, AirdishStraus, Aitias, Alansohn, Ale jrb, Alexius08,
[email protected], All Is One, Alvis, Andre Engels, AndreasJS, Andres, Anonymous editor, Anoop.m, Antandrus, Anthony Appleyard, Antidisestablishmentarinism, Any Moose, Anárion, Apostrophe, Aranherunar, Archfiendweazal, Archimerged, Ari639, Arjun01, Art LaPella, Atlant, Auximines, Awolf002, Azn king28, B.d.mills, BRG, BW, Bachrach44, Badocter, Beardo, Beetstra, Bergsten, Bettia, Bjweeks, Bkonrad, BlueEarth, Bobblewik, Bobbo, Bobet, Bobo192, Bomac, Brian0918, Brockert, Brooknet, Bryan Derksen, BunsenH, CTZMSC3, CYD, Caltas, Can't sleep, clown will eat me, CaptainVindaloo, Carnildo, Casper2k3, Catgut, Ceranthor, ChicXulub, Chris Dybala, Christian List, Chuto, Closedmouth, Computermix, ConradPino, Conversion script, Corpx, Crazy Boris with a red beard, Cremepuff222, Cryptic C62, Cureden, Curps, Cyclist 1977, Dabomb87, Daemon Reborn, Dannyc77, Darrien, Dave McKee, David Latapie, Ddcampayo, DeadShort, Deglr6328, Delirium, DerHexer, Derek.cashman, Dgrant, Dhall12345, Dillona, Dirac66, Discospinster, Donald Albury, Donarreiskoffer, Doulos Christos, DragonflySixtyseven, Drini, Dschor, Dsmouse, Dspradau, E Wing, EL Willy, Ec5618, Edgar181, Edsanville, Eeekster, El C, Elaragirl, Elemesh, Elkman, Emmy.rocks, Emperorbma, Eng02019, Enok Walker, Epbr123, Erebus Morgaine, Eric Shalov, Eric119, Erik Zachte, Evil saltine, Extraordinary, FaerieInGrey, Farosdaughter, Femto, Flauto Dolce, Flewis, Fonzy, Foobar, Frankenpuppy, Freestyle-69, FreplySpang, Fresheneesz, Funky Monkey, GT5162, Garden, Gcsuchemistry, Gene Hobbs, Gene Nygaard, Geoffr, Giftlite, Gilliam, Gman124, Gorm, Goudzovski, Graydonflatt, Grendelkhan, Grin, Gtdp, Gtstricky, Gubbubu, Gurch, Gwernol, Hadal, Haham hanuka, Hak-kâ-ngìn, Hawaii fiveoh, Hda3ku, Heron, Humormekil, Hut 8.5, Hvn0413, Hydrogen Iodide, II MusLiM HyBRiD II, IRP, Iago4096, Icairns, Iluvcapra, Indon, Indosauros, Insanity Incarnate, Itub, Ixfd64, J.delanoy, JaGa, Jacek Kendysz, Jagun, Jake Nelson, Jaknouse, Janke, Jaraalbe, Jayden54, Jehan UK, Jespinos, JetLover, Jimp, Jj137, Joanjoc, JoanneB, JoeSmack, John, John Holmes II, Jokah49, Jonadab, Jose77, Jossi, Juliancolton, Junglecat, Karl-Henner, Karlhahn, Kaszeta, King of Hearts, Kingpin13, Klingoncowboy4, Kmweber, KnowledgeOfSelf, Kostisl, Kowey, Kozuch, Krich, Ksbrown, Kuithei, Kurykh, Kwamikagami, Lamanatasa, LeaveSleaves, Leonard^Bloom, Lexor, Lightmouse, Ling.Nut, Livajo, Llywelyn, LouScheffer, Lradrama, LuigiManiac, Lyoko is Cool, M100, M412k, MBisanz, MZMcBride, Magu2k, Mandarax, Marek69, MarkS, Markhurd, Marnanel, Materialscientist, Matt Deres, Mav, MayaSimFan, Merlincooper, Merovingian, Messs17, Mgimpel, Mike Rosoft, Minesweeper, Minnesota1, Mistyfrog, Misza13, Mitsuhama, Mossman93, Mr.Z-man, Mr45acp, Mulad, Mygerardromance, N328KF, Nakon, Nancy, Naraht, Natalie Erin, NawlinWiki, Neonfrank, Nergaal, Neverquick, NewEnglandYankee, Nick C, Night Gyr, Nihiltres, No1lakersfan, Novangelis, Nsaa, Number 57, Opabinia regalis, Ossmann, Oxymoron83, PP Jewel, Pascal.Tesson, Passw0rd, Patrick Berry, Paul August, Peripitus, PeterJeremy, Pewwer42, Pfunk42, PhilKnight, Philip Trueman, PhilipO, Piano non troppo, Pierre.laloë, Plantsurfer, Plasmic Physics, Plasticup, Plinkit, Poltair, Poolkris, Possum, PrimaryDragon, Psyche825, Ptdecker, Quadell, Quintote, Quistnix, RAM, RTC, Rdsmith4, RedWolf, Redux, Remember, Res2216firestar, RexNL, Richnotts, Rifleman 82, Riflewhipu, Rodhullandemu, Romanm, RyanB88, RyanCross, Ryanh1994, S lijin, S0me l0ser, SEWilco, SJP, Sam Korn, Saperaud, Sbharris, SchfiftyThree, Schneelocke, Seddon, Sengkang, Sepharious, Sfgagnon, Shadowdude77, Shaun F, Shellreef, Shenme, SimonP, Sinblox, Sionus, Sjö, Sl, Solipsist, SouthernSunShine, Squids and Chips, Staphylococcus, StaticVision, SteinbDJ, Stephenb, Steve Crossin, Stifynsemons, Stone, Suisui, Sultanofsurreal, Sunborn, Tapir Terrific, Tarret, Techman224, Tham153, TheCatalyst31, TheDJ, TheKMan, TheNewPhobia, Thefirstgolfoffer, Thej42, Thief111362, Thricecube, Tide rolls, TigerShark, Tiptoety, Titoxd, Tlim7882, Tombomp, Tomwzhang95, TonyW, Treelo, Trevor MacInnis, Trigger820, Tsogo3, TutterMouse, Until It Sleeps, Urhixidur, Vancouverguy, Versus22, Video game fan11, Vmatikov, Voyagerfan5761, Vsmith, Vssun, Vuong Ngan Ha, WLU, Wang lvan, Watch37264, Weeliljimmy, Whitepaw, Whosasking, Wikidougg, Wilbern Cobb, Willking1979, Wimt, WinterSpw, Wolfkeeper, Woodster93, X201, Yamamoto Ichiro, Yath, Youandme, Yyy, Zach4636, Zm786, 1017 anonymous edits
Image Sources, Licenses and Contributors file:cubic-face-centered.svg Source: http://en.wikipedia.org/w/index.php?title=File:Cubic-face-centered.svg License: GNU Free Documentation License Contributors: User:Stannered file:Electron shell 018 Argon.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_018_Argon.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 Image:Isolation of Argon.png Source: http://en.wikipedia.org/w/index.php?title=File:Isolation_of_Argon.png License: Public Domain Contributors: Encyclopædia Britannica Image:Argon ice 1.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Argon_ice_1.jpg License: unknown Contributors: Brian0918, Er Komandante, 2 anonymous edits Image:Argon.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Argon.jpg License: Attribution Contributors: William Viker File:CsCrystals.JPG Source: http://en.wikipedia.org/w/index.php?title=File:CsCrystals.JPG License: unknown Contributors: User:Dnn87 Image:Glovebox.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Glovebox.jpg License: GNU Free Documentation License Contributors: Polimerek Image:ArTube.jpg Source: http://en.wikipedia.org/w/index.php?title=File:ArTube.jpg License: unknown Contributors: User:Pslawinski
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10
Potassium
1
Potassium argon ← potassium → calciumNa ↑ K ↓ Rb
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
19K Periodic table
Appearance silvery white
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointTriple pointHeat of fusionHeat of vaporizationSpecific heat capacity Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 3052 kJ·mol−1 3rd: 4420 kJ·mol−1Atomic radiusCovalent radiusVan der Waals radius Miscellaneous Crystal structureMagnetic orderingThermal conductivityThermal
Potassium
2
expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusMohs hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of potassium iso
N.A.
half-life
39
93.26%
39
40
0.012%
1.248(3)×109 y
K K
41
K
6.73%
DM
DE (MeV)
DP
K is stable with 20 neutron β−
1.311
40
ε
1.505
40
β+
1.505
40
Ca Ar Ar
41
K is stable with 22 neutron
potassium, K, 19 alkali metal1, 4, s39.0983(1) g·mol−1 [Ar] 4s1 2, 8, 8, 1 (Image) solid 0.89 g·cm−3 0.828 g·cm−3 336.53 K,63.38 °C,146.08 °F 1032 K,759 °C,1398 °F −1 −1 −1 −1 336.35 K (63°C), kPa 2.33 kJ·mol 76.9 kJ·mol (25 °C) 29.6 J·mol ·K 1 (strongly basic oxide) 0.82 (Pauling scale) 1st: 418.8 kJ·mol−1227 pm203±12 pm 275 pm body-centered cubic paramagnetic (300 K) 102.5 W·m−1·K−1 (25 °C) 83.3 µm·m−1·K−1 (20 °C) 2000 m/s 3.53 GPa 1.3 GPa 3.1 GPa 0.4 0.363 MPa 7440-09-7 Potassium (pronounced /pɵˈtæsiəm/) is the chemical element with the symbol K (Latin: kalium, from Arabic: هَيْلَقلا al-qalyah “plant ashes”, cf. Alkali from the same root), atomic number 19, and atomic mass 39.0983. Potassium was first isolated from potash. Elemental potassium is a soft silvery-white metallic alkali metal that oxidizes rapidly in air and is very reactive with water, generating sufficient heat to ignite the evolved hydrogen. Potassium in nature occurs only as ionic salt. As such, it is found dissolved in seawater, and as part of many minerals. Potassium ion is necessary for the function of all living cells, and is thus present in all plant and animal tissues. It is found in especially high concentrations in plant cells, and in a mixed diet, it is most highly concentrated in fruits. In many respects, potassium and sodium are chemically similar, although they have very different functions in organisms in general, and in animal cells in particular.
Occurrence Elemental potassium does not occur in nature because it reacts violently with water.[1] As various compounds, potassium makes up about 1.5% of the weight of the Earth's crust and is the seventh most abundant element.[1] As it is very electropositive, potassium metal is difficult to obtain from its minerals.
History of the free element Potassium in feldspar
Elemental potassium was not known in Roman times, and its names are not Classical Latin but rather neo-Latin. The name kalium was taken from the word
"alkali", which came from Arabic al qalīy = "the calcined ashes". The name potassium was made from the word "potash", which is English, and originally meant an alkali extracted in
Potassium
3
a pot from the ash of burnt wood or tree leaves. The structure of potash was not then known, but is now understood to be mostly potassium carbonate. By heating, the carbonate could be freed of carbon dioxide, leaving "caustic potash", so called because it caused chemical burns in contact with human tissue. Potassium metal was discovered in 1807 in England by Sir Humphry Davy, who derived it from caustic potash (KOH), by the use of electrolysis of the molten salt with the newly discovered voltaic pile. Before the 18th century, no distinction was made between potassium and sodium. Potassium was the first metal that was isolated by electrolysis.[2] Davy extracted sodium by a similar technique, demonstrating the elements to be different.[3]
Production Pure potassium metal may be isolated by electrolysis of its hydroxide in a process that has changed little since Davy.[1] Thermal methods also are employed in potassium production, using potassium chloride Potassium salts such as carnallite, langbeinite, polyhalite, and sylvite form extensive deposits in ancient lake and seabeds, making extraction of potassium salts in these environments commercially viable. The principal source of potassium, potash, is mined in Saskatchewan, California, Germany, New Mexico, Utah, and in other places around the world. It is also found abundantly in the Dead Sea. Three thousand feet below the surface of Saskatchewan are large deposits of potash which are important sources of this element and its salts, with several large mines in operation since the 1960s. Saskatchewan pioneered the use of freezing of wet sands (the Blairmore formation) in order to drive mine shafts through them. The main mining company is the Potash Corporation of Saskatchewan. The oceans are another source of potassium, but the quantity present in a given volume of seawater is relatively low compared with sodium.[4] [5]
Isotopes 39
40
There are 24 known isotopes of potassium. Three isotopes occur naturally: K (93.3%), K (0.0117%) and 41K (6.7%). Naturally occurring 40K decays to stable 40Ar (11.2%) by electron capture and by positron emission, and decays to stable 40Ca (88.8%) by beta decay; 40K has a half-life of 1.250×109 years. The decay of 40K to 40Ar enables a commonly used method for dating rocks. The conventional K-Ar dating method depends on the assumption that the rocks contained no argon at the time of formation and that all the subsequent radiogenic argon (i.e., 40Ar) was quantitatively retained. Minerals are dated by measurement of the concentration of potassium and the amount of radiogenic 40Ar that has accumulated. The minerals that are best suited for dating include biotite, muscovite, plutonic/high grade metamorphic hornblende, and volcanic feldspar; whole rock samples from volcanic flows and shallow instrusives can also be dated if they are unaltered. Outside of dating, potassium isotopes have been used extensively as tracers in studies of weathering. They have also been used for nutrient cycling studies because potassium is a macronutrient required for life. 40
K occurs in natural potassium (and thus in some commercial salt substitutes) in sufficient quantity that large bags of those substitutes can be used as a radioactive source for classroom demonstrations. In healthy animals and people, 40K represents the largest source
Potassium of radioactivity, greater even than of 40K decay per second.[6]
4 14
C. In a human body of 70 kg mass, about 4,400 nuclei
The activity of natural potassium is 31 Bq/g.
Properties Physical Potassium is the second least dense metal; only lithium is less dense. It is a soft, low-melting solid that can easily be cut with a knife. Freshly cut potassium is silvery in appearance, but in air it begins to tarnish toward grey immediately.[1] In a flame test, potassium and its compounds emit a pale violet color, which may be masked by the strong yellow emission of sodium if it is also present. Cobalt glass can be used to filter out the yellow sodium color.[7] Potassium concentration in solution is commonly determined by flame photometry, atomic absorption spectrophotometry, inductively coupled plasma, or ion selective electrodes.
Chemical Potassium must be protected from air for storage to prevent disintegration of the metal from oxide and hydroxide corrosion. Often samples are maintained under a hydrocarbon medium which does not react with alkali metals, such as mineral oil or kerosene. Like the other alkali metals, potassium reacts violently with water, producing hydrogen. The reaction is notably more violent than that of lithium or sodium with water, and is sufficiently exothermic that the evolved hydrogen gas ignites.
The flame-test color for potassium
2K(s) + 2H2O(l) → H2(g) + 2KOH(aq) Because potassium reacts quickly with even traces of water, and its reaction products are nonvolatile, it is sometimes used alone, or as NaK (an alloy with sodium which is liquid at room temperature) to dry solvents prior to distillation. In this role, it serves as a potent desiccant. Potassium hydroxide reacts strongly with carbon dioxide to produce potassium carbonate, and is used to remove traces of CO2 from air. Potassium compounds generally have excellent water solubility, due to the high hydration energy of the K+ ion. The potassium ion is colorless in water. Methods of separating potassium by precipitation, sometimes used for gravimetric analysis, include the use of sodium tetraphenylborate, hexachloroplatinic acid, and sodium cobaltinitrite
Potassium
Potassium cations in the body Biochemical function Potassium cations are important in neuron (brain and nerve) function, and in influencing osmotic balance between cells and the interstitial fluid, with their distribution mediated in all animals (but not in all plants) by the so-called Na+/K+-ATPase pump.[8] Potassium may be detected by taste because it triggers three of the five types of taste sensations, according to concentration. Dilute solutions of potassium ion taste sweet (allowing moderate concentrations in milk and juices), while higher concentrations become increasingly bitter/alkaline, and finally also salty to the taste. The combined bitterness and saltiness of high potassium content solutions makes high-dose potassium supplementation by liquid drinks a palatability challenge.[9]
Membrane polarization Potassium is also important in allowing muscle contraction and the sending of all nerve impulses in animals through action potentials. By nature of their electrostatic and chemical properties, K+ ions are larger than Na+ ions, and ion channels and pumps in cell membranes can distinguish between the two types of ions, actively pumping or passively allowing one of the two ions to pass, while blocking the other. [10] A shortage of potassium in body fluids may cause a potentially fatal condition known as hypokalemia, typically resulting from diarrhea, increased diuresis and vomiting. Deficiency symptoms include muscle weakness, paralytic ileus, ECG abnormalities, decreased reflex response and in severe cases respiratory paralysis, alkalosis and cardiac arrhythmia.
Filtration and excretion Potassium is an essential mineral micronutrient in human nutrition; it is the major cation (positive ion) inside animal cells, and it is thus important in maintaining fluid and electrolyte balance in the body. Sodium makes up most of the cations of blood plasma at a reference range of about 145 milliequivalents per liter (3345 milligrams) and potassium makes up most of the cell fluid cations at about 150 milliequivalents per liter (4800 milligrams). Plasma is filtered through the glomerulus of the kidneys in enormous amounts, about 180 liters per day.[11] Thus 602,000 milligrams of sodium and 33,000 milligrams of potassium are filtered each day. All but the 1000-10,000 milligrams of sodium and the 1000-4000 milligrams of potassium likely to be in the diet must be reabsorbed. Sodium must be reabsorbed in such a way as to keep the blood volume exactly right and the osmotic pressure correct; potassium must be reabsorbed in such a way as to keep serum concentration as close as possible to 4.8 milliequivalents (about 190 milligrams) per liter.[12] Sodium pumps in the kidneys must always operate to conserve sodium. Potassium must sometimes be conserved also, but since the amount of potassium in the blood plasma is very small and the pool of potassium in the cells is about thirty times as large, the situation is not so critical for potassium. Since potassium is moved passively[13] [14] in counter flow to sodium in response to an apparent (but not actual) Donnan equilibrium,[15] the urine can never sink below the concentration of potassium in serum except sometimes by actively excreting water at the end of the processing. Potassium is secreted twice and reabsorbed three times before the urine reaches the collecting tubules.[16] At that point, it usually has about the same potassium concentration as plasma. If potassium were removed
5
Potassium from the diet, there would remain a minimum obligatory kidney excretion of about 200 mg per day when the serum declines to 3.0-3.5 milliequivalents per liter in about one week,[17] and can never be cut off completely. Because it cannot be cut off completely, death will result when the whole body potassium declines to the vicinity of one-half full capacity. At the end of the processing, potassium is secreted one more time if the serum levels are too high.
Reference ranges for blood tests, showing blood content of potassium (3.6 to 5.2 mmol/l) in blue in right part of the spectrum.
The potassium moves passively through pores in the cell wall. When ions move through pumps there is a gate in the pumps on either side of the cell wall and only one gate can be open at once. As a result, 100 ions are forced through per second. Pores have only one gate, and there only one kind of ion can stream through, at 10 million to 100 million ions per second.[18] The pores require calcium in order to open[19] although it is thought that the calcium works in reverse by blocking at least one of the pores.[20] Carbonyl groups inside the pore on the amino acids mimics the water hydration that takes place in water solution[21] by the nature of the electrostatic charges on four carbonyl groups inside the pore.[22]
Potassium in the diet and by supplement Adequate intake A potassium intake sufficient to support life can generally be guaranteed by eating a variety of foods, especially plant foods. Clear cases of potassium deficiency (as defined by symptoms, signs and a below-normal blood level of the element) are rare in healthy individuals eating a balanced diet. Foods with high sources of potassium include orange juice, potatoes, bananas, avocados, tomatoes, broccoli, soybeans, brown rice, garlic and apricots, although it is also common in most fruits, vegetables and meats.[23] Optimal intake Epidemiological studies and studies in animals subject to hypertension indicate that diets high in potassium can reduce the risk of hypertension and possibly stroke (by a mechanism independent of blood pressure), and a potassium deficiency combined with an inadequate thiamine intake has produced heart disease in rats.[24] With these findings, the question of what is the intake of potassium consistent with optimal health, is debated. For example, the 2004 guidelines of the Institute of Medicine specify a DRI of 4,000 mg of potassium (100 mEq), though most Americans consume only half that amount per day, which would make them formally deficient as regards this particular recommendation.[25] Similarly, in the European Union, particularly in Germany and Italy, insufficient potassium intake is somewhat common.[26]
6
Potassium
7
Medical supplementation and disease Supplements of potassium in medicine are most widely used in conjunction with loop diuretics and thiazides, classes of diuretics which rid the body of sodium and water, but have the side effect of also causing potassium loss in urine. A variety of medical and non-medical supplements are available. Potassium salts such as potassium chloride may be dissolved in water, but the salty/bitter taste of high concentrations of potassium ion make palatable high concentration liquid supplements difficult to formulate.[9] Typical medical supplemental doses range from 10 milliequivalents (400 mg, about equal to a cup of milk or 6 oz. of orange juice) to 20 milliequivalents (800 mg) per dose. Potassium salts are also available in tablets or capsules, which for therapeutic purposes are formulated to allow potassium to leach slowly out of a matrix, since very high concentrations of potassium ion (which might occur next to a solid tablet of potassium chloride) can kill tissue, and cause injury to the gastric or intestinal mucosa. For this reason, non prescription supplement potassium pills are limited by law in the U.S. to only 99 mg of potassium. Individuals suffering from kidney diseases may suffer adverse health effects from consuming large quantities of dietary potassium. End stage renal failure patients undergoing therapy by renal dialysis must observe strict dietary limits on potassium intake, since the kidneys control potassium excretion, and buildup of blood concentrations of potassium (hyperkalemia) may trigger fatal cardiac arrhythmia.
Applications About 93% of the world potassium production was consumed by the fertilizer industry.[5]
Biological applications Potassium ions are an essential component of plant nutrition and are found in most soil types. Its primary use in agriculture, horticulture and hydroponic culture is as a fertilizer as the chloride (KCl), sulfate (K2SO4) or nitrate (KNO3). In animal cells, potassium ions are vital to keeping cells alive (see Na-K pump). In the form of potassium chloride, it is used to stop the heart, e.g. in cardiac surgery and in a solution used in executions by lethal injection.
Food applications
Potassium and Magnesium sulfate fertilizer
Potassium ion is a nutrient necessary for human life and health. Potassium chloride is used as a substitute for table salt by those seeking to reduce sodium intake so as to control hypertension. The USDA lists tomato paste, orange juice, beet greens, white beans, potatoes, bananas and many other good dietary sources of potassium, ranked according to potassium content per measure shown.[27] Potassium sodium tartrate, or Rochelle salt (KNaC4H4O6) is the main constituent of baking powder. Potassium bromate (KBrO3) is a strong oxidiser, used as a flour improver (E924) to
Potassium improve dough strength and rise height. The sulfite compound, potassium bisulfite (KHSO3) is used as a food preservative, for example in wine and beer-making (but not in meats). It is also used to bleach textiles and straw, and in the tanning of leathers.
Industrial applications Potassium vapor is used in several types of magnetometers. An alloy of sodium and potassium, NaK (usually pronounced "nack"), that is liquid at room temperature, is used as a heat-transfer medium. It can also be used as a desiccant for producing dry and air-free solvents. Potassium metal reacts vigorously with all of the halogens to form the corresponding potassium halides, which are white, water-soluble salts with cubic crystal morphology. Potassium bromide (KBr), potassium iodide (KI) and potassium chloride (KCl) are used in photographic emulsion to make the corresponding photosensitive silver halides. Potassium hydroxide KOH is a strong base, used in industry to neutralize strong and weak acids and thereby finding uses in pH control and in the manufacture of potassium salts. Potassium hydroxide is also used to saponify fats and oils and in hydrolysis reactions, for example of esters and in industrial cleaners. Potassium nitrate KNO3 or saltpeter is obtained from natural sources such as guano and evaporites or manufactured by the Haber process and is the oxidant in gunpowder (black powder) and an important agricultural fertilizer. Potassium cyanide KCN is used industrially to dissolve copper and precious metals particularly silver and gold by forming complexes; applications include gold mining, electroplating and electroforming of these metals. It is also used in organic synthesis to make nitriles. Potassium carbonate K2CO3, also known as potash, is used in the manufacture of glass and soap and as a mild desiccant. Potassium chromate (K2CrO4) is used in inks, dyes, and stains (bright yellowish-red colour), in explosives and fireworks, in safety matches, in the tanning of leather and in fly paper. Potassium fluorosilicate (K2SiF6) is used in specialized glasses, ceramics, and enamels. Potassium sodium tartrate, or Rochelle salt (KNaC4H4O6) is used in the silvering of mirrors. The superoxide KO2 is an orange-colored solid used as a portable source of oxygen and as a carbon dioxide absorber. It is useful in portable respiration systems. It is widely used in submarines and spacecraft as it takes less volume than O2 (g). 4KO2 + 2CO2 --> 2K2CO3 + 3O2 Potassium chlorate KClO3 is a strong oxidant, used in percussion caps and safety matches and in agriculture as a weedkiller. Glass may be treated with molten potassium nitrate KNO3 to make toughened glass, which is much stronger than regular glass.
Precautions Potassium reacts very violently with water producing hydrogen gas which then usually catches fire. Potassium is usually kept under a hydrocarbon oil such as mineral oil or kerosene to stop the metal from reacting with water vapour present in the air. Unlike lithium and sodium, however, potassium should not be stored under oil indefinitely. If stored longer than 6 months to a year, dangerous shock-sensitive peroxides can form on the metal and under the lid of the container, which can detonate upon opening. It is recommended that potassium, rubidium or caesium not be stored for longer than three
8
Potassium
9
months unless stored in an inert (oxygen free) atmosphere, or under vacuum.[28] As potassium reacts with water to produce highly flammable hydrogen gas, a potassium fire is only exacerbated by the addition of water, and only a few dry chemicals are effective for putting out such a fire (see the precaution section in sodium). Potassium also produces potassium hydroxide (KOH) in the reaction with water. Potassium hydroxide is a strong alkali and so is a caustic hazard, causing burns. Due to the highly reactive nature of potassium metal, it must be handled with great care, with full skin and eye protection being used and preferably an explosive resistant barrier between the user and the potassium.
See also • Potassium compounds • Potassium in biology
External links • WebElements.com – Potassium
[29]
References [1] Mark Winter. " Potassium: Key Information (http:/ / www. webelements. com/ webelements/ elements/ text/ K/ key. html)". Webelements. . [2] Enghag, P. (2004). "11. Sodium and Potassium". Encyclopedia of the elements. Wiley-VCH Weinheim. ISBN 3527306668. [3] Davy, Humphry (1808). " On some new Phenomena of Chemical Changes produced by Electricity, particularly the Decomposition of the fixed Alkalies, and the Exhibition of the new Substances, which constitute their Bases (http:/ / books. google. com/ books?id=Kg9GAAAAMAAJ)". Philosophical Transactions of the Royal Society of London 98: 1–45. doi: 10.1098/rstl.1808.0001 (http:/ / dx. doi. org/ 10. 1098/ rstl. 1808. 0001). . [4] Ober, Joyce A.. " Mineral Commodity Summaries 2008:Potash (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ potash/ mcs-2008-potas. pdf)". United States Geological Survey. . Retrieved 2008-11-20. [5] Ober, Joyce A.. " Mineral Yearbook 2006:Potash (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ potash/ myb1-2006-potas. pdf)". United States Geological Survey. . Retrieved 2008-11-20. [6] " background radiation - potassium-40 - γ radiation (http:/ / www. fas. harvard. edu/ ~scdiroff/ lds/ QuantumRelativity/ RadioactiveHumanBody/ RadioactiveHumanBody. html)". . [7] Anne Marie Helmenstine. " Qualitative Analysis - Flame Tests (http:/ / chemistry. about. com/ library/ weekly/ aa110401a. htm)". About.com. . [8] Campbell, Neil (1987). Biology. Menlo Park, Calif.: Benjamin/Cummings Pub. Co.. pp. 795. ISBN 0-8053-1840-2. [9] " Potassium Without the Taste (http:/ / www. foodnavigator. com/ Science-Nutrition/ Potassium-without-the-taste)". . Retrieved Feb 14, 2009. [10] Lockless SW, Zhou M, MacKinnon R.. " Structural and thermodynamic properties of selective ion binding in a K+ channel (http:/ / www. ncbi. nlm. nih. gov/ pubmed/ 17472437)". Laboratory of Molecular Neurobiology and Biophysics, Rockefeller University. . Retrieved 2008-03-08. [11] Potts, W.T.W.; Parry, G. (1964). Osmotic and ionic regulation in animals. Pergamon Press. [12] Lans HS, Stein IF, Meyer KA (1952). "The relation of serum potassium to erythrocyte potassium in normal subjects and patients with potassium deficiency". Am. J. Med. Sci. 223 (1): 65–74. doi: 10.1097/00000441-195201000-00011 (http:/ / dx. doi. org/ 10. 1097/ 00000441-195201000-00011). PMID 14902792. [13] Bennett CM, Brenner BM, Berliner RW (1968). "Micropuncture study of nephron function in the rhesus monkey". J Clin Invest 47 (1): 203–216. PMID 16695942. [14] Solomon AK (1962). "Pumps in the living cell". Sci. Am. 207: 100–8. PMID 13914986. [15] Kernan, Roderick P. (1980). Cell potassium (Transport in the life sciences). New York: Wiley. ISBN 0471048062.; p. 40 & 48.
Potassium [16] Wright FS (1977). "Sites and mechanisms of potassium transport along the renal tubule". Kidney Int. 11 (6): 415–32. doi: 10.1038/ki.1977.60 (http:/ / dx. doi. org/ 10. 1038/ ki. 1977. 60). PMID 875263. [17] Squires RD, Huth EJ (1959). "Experimental potassium depletion in normal human subjects. I. Relation of ionic intakes to the renal conservation of potassium". J. Clin. Invest. 38 (7): 1134–48. doi: 10.1172/JCI103890 (http:/ / dx. doi. org/ 10. 1172/ JCI103890). PMID 13664789. [18] Gadsby DC (2004). "Ion transport: spot the difference". Nature 427 (6977): 795–7. doi: 10.1038/427795a (http:/ / dx. doi. org/ 10. 1038/ 427795a). PMID 14985745.; for a diagram of the potassium pores are viewed, see Miller C (2001). "See potassium run". Nature 414 (6859): 23–4. doi: 10.1038/35102126 (http:/ / dx. doi. org/ 10. 1038/ 35102126). PMID 11689922. [19] Jiang Y, Lee A, Chen J, Cadene M, Chait BT, MacKinnon R (2002). "Crystal structure and mechanism of a calcium-gated potassium channel". Nature 417 (6888): 515–22. doi: 10.1038/417515a (http:/ / dx. doi. org/ 10. 1038/ 417515a). PMID 12037559. [20] Shi N, Ye S, Alam A, Chen L, Jiang Y (2006). "Atomic structure of a Na+- and K+-conducting channel". Nature 440 (7083): 570–4. doi: 10.1038/nature04508 (http:/ / dx. doi. org/ 10. 1038/ nature04508). PMID 16467789.; includes a detailed picture of atoms in the pump. [21] Zhou Y, Morais-Cabral JH, Kaufman A, MacKinnon R (2001). "Chemistry of ion coordination and hydration revealed by a K+ channel-Fab complex at 2.0 A resolution". Nature 414 (6859): 43–8. doi: 10.1038/35102009 (http:/ / dx. doi. org/ 10. 1038/ 35102009). PMID 11689936. [22] Noskov SY, Bernèche S, Roux B (2004). "Control of ion selectivity in potassium channels by electrostatic and dynamic properties of carbonyl ligands". Nature 431 (7010): 830–4. doi: 10.1038/nature02943 (http:/ / dx. doi. org/ 10. 1038/ nature02943). PMID 15483608. [23] " Potassium Content of Food and Drink (http:/ / www. pamf. org/ patients/ pdf/ potassium_count. pdf)". . Retrieved 2008-09-18. [24] Folis, R.H. (1942). "Myocardial Necrosis in Rats on a Potassium Low Diet Prevented by Thiamine Deficiency". Bull. Johns-Hopkins Hospital 71: 235. [25] Grim CE, Luft FC, Miller JZ, et al. (1980). "Racial differences in blood pressure in Evans County, Georgia: relationship to sodium and potassium intake and plasma renin activity". J Chronic Dis 33 (2): 87–94. doi: 10.1016/0021-9681(80)90032-6 (http:/ / dx. doi. org/ 10. 1016/ 0021-9681(80)90032-6). PMID 6986391. [26] Karger, S. (2004). " Energy and nutrient intake in the European Union (http:/ / content. karger. com/ ProdukteDB/ produkte. asp?Aktion=ShowPDF& ProduktNr=223977& Ausgabe=230671& ArtikelNr=83312& filename=83312. pdf)" (pdf). Ann Nutr Metab 48 (2 (suppl)): 1–16. . [27] )Potassium / K (mg.) Content of Selected Foods per Common Measure, sorted by nutrient content | USDA National Nutrient Database for Standard Reference, Release 20 http:/ / www. nal. usda. gov/ fnic/ foodcomp/ Data/ SR20/ nutrlist/ sr20w306. pdf [28] Thomas K. Wray. " DANGER: PEROXIDIZABLE CHEMICALS (http:/ / www. ncsu. edu/ ehs/ www99/ right/ handsMan/ lab/ Peroxide. pdf)". Environmental Health & Public Safety (North Carolina State University). . [29] http:/ / www. webelements. com/ webelements/ elements/ text/ K/ index. html
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Article Sources and Contributors
Article Sources and Contributors Potassium Source: http://en.wikipedia.org/w/index.php?oldid=307655059 Contributors: (jarbarf), 2D, 417patriot, 64.174.7.xxx, 65.68.87.xxx, A2Kafir, ABF, Acebulf, Acer77c, Adashiel, Addshore, Aeluwas, Agemegos, Aglomax, Ahoerstemeier, Aitias, Aksi great, Alansohn, Aleron235, Alexalexalex39, AlexiusHoratius, Aliphatic Hydrocarbon, Almost222, Amaggi1, Ambar, AndreasJS, Andres, Andrewa, Andrewpmk, Andypham3000, Anlace, AnnaFrance, Anonymous Dissident, Antandrus, Anthony Appleyard, Archimerged, ArchonMagnus, Astgtciv, AuburnPilot, AussieWiki2006, BTLizard, BaysiderGene, Beetstra, Big Bird, Biochemnick, BlindEagle, BlueEarth, Bobet, Bobo192, Bogey97, Bomac, Bonorox, BorgHunter, Bornhj, Bryan Derksen, Buchanan-Hermit, Bunbury18, CTZMSC3, CWii, CYD, CactusWriter, Calor, Can't sleep, clown will eat me, CapitalR, Capricorn42, CaptainCarrot, Carnildo, Chameleon, ChemNerd, Chmod007, Chris 73, Christopher Parham, Chromaticity, Chubleathem, Clocktopus, Closedmouth, Colbuckshot, Cometstyles, Connorhalsell, Conversion script, Coppertwig, Courseiam, Crabula, Crazysane, Creidieki, Crystal whacker, Cspurrier, Cunado19, Cureden, Cyrius, D, DCEdwards1966, DMacks, DVD R W, Da monster under your bed, Dajwilkinson, Damburger, DanArcari, Daniel.Cardenas, DanielCD, Darrien, David Latapie, David Schaich, David044058, DavidOaks, Db099221, Ddillon, Dead3y3, DeadEyeArrow, Deanos, Delta G, Deltabeignet, Denelson83, DerHexer, Dferg, Digresser, Discospinster, Djayjp, Dlohcierekim, Dnn87, Donarreiskoffer, Doodle77, Dr.Soft, DrBob, Dreadstar, Dspradau, Duhhhhhh, Dukeofwulf, Duncan Smikle, Dysepsion, E0steven, EL Willy, EamonnPKeane, Eclecticology, EcwNick, Edgar181, Edwy, Eequor, Ehrenkater, El C, Eldin raigmore, Elkman, Emfetz, Emperorbma, Enviroboy, Epbr123, Eric Kvaalen, Eric-Wester, Erik Zachte, Erpacx14, Eve Hall, Everyking, Ewlyahoocom, Faithlessthewonderboy, Fan-1967, Femto, Fieldday-sunday, Fitzy93, Flewis, Floydrox, Flyguy649, Focusss, Fonzy, Frankenpuppy, FreplySpang, Friedrich V. 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11
Calcium
1
Calcium potassium ← calcium → scandiumMg ↑ Ca ↓ Sr
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
20Ca Periodic table
Appearance Dull grey, silver
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
864
956
1071
1227
1443
1755
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1145.4 kJ·mol−1 3rd: 4912.4 kJ·mol−1Atomic radiusCovalent radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal
Calcium
2
conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusPoisson ratioMohs hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of calcium iso
N.A.
half-life
40
96.941%
40
41
syn
1.03×105 y
42
0.647%
42
43
0.135%
43
44
2.086%
44
45
syn
162.7 d
46
0.004%
47
syn
Ca Ca Ca Ca Ca Ca Ca Ca
48
Ca
0.187%
DM
DE (MeV)
DP
Ca is stable with 20 neutron ε
-
41
β−
0.258
45
>2.8×1015 y
β−β−
?
46
4.536 d
β−
0.694, 1.99
47
γ
1.297
-
β−β−
?
48
K
Ca is stable with 22 neutron Ca is stable with 23 neutron Ca is stable with 24 neutron
>4×1019 y
Sc Ti Sc
Ti
calcium, Ca, 20 alkaline earth metal2, 4, s40.078(4) g·mol−1 [Ar] 4s2 2, 8, 8, 2 (Image) solid 1.55 g·cm−3 1.378 g·cm−3 1115 K,842 °C,1548 °F 1757 K,1484 °C,2703 °F 8.54 kJ·mol−1 154.7 kJ·mol−1 (25 °C) 25.929 J·mol−1·K−1 2 (strongly basic oxide) 1.00 (Pauling scale) 1st: 589.8 kJ·mol−1197 pm176±10 pm face-centered cubic diamagnetic (20 °C) 33.6 nΩ·m (300 K) 201 W·m−1·K−1 (25 °C) 22.3 µm·m−1·K−1 (20 °C) 3810 m/s 20 GPa 7.4 GPa 17 GPa 0.31 1.75 167 MPa 7440-70-2 Calcium (pronounced /ˈkælsiəm/) is the chemical element with the symbol Ca and atomic number 20. It has an atomic mass of 40.078 amu. Calcium is a soft grey alkaline earth metal, and is the fifth most abundant element by mass in the Earth's crust. Calcium is also the fifth most abundant dissolved ion in seawater by both molarity and mass, after sodium, chloride, magnesium, and sulfate.[1] Calcium is essential for living organisms, particularly in cell physiology, where movement of the calcium ion Ca2+ into and out of the cytoplasm functions as a signal for many cellular processes. As a major material used in mineralization of bones and shells, calcium is the most abundant metal by mass in many animals.
Calcium
Notable characteristics Chemically calcium is reactive and soft for a metal (though harder than lead, it can be cut with a knife with difficulty). It is a silvery metallic element that must be extracted by electrolysis from a fused salt like calcium chloride.[2] Once produced, it rapidly forms a grey-white oxide and nitride coating when exposed to air. It is somewhat difficult to ignite, unlike magnesium, but when lit, the metal burns in air with a brilliant high-intensity red light. Calcium metal reacts with Calcium carbonate wetted with water, evolving hydrogen gas at a rate rapid enough to hydrochloric acid (thus forming CaCl2) be noticeable, but not fast enough at room temperature held at a flame and showing to generate much heat. In powdered form, however, the red-orange flame color of Ca. reaction with water is extremely rapid, as the increased surface area of the powder accelerates the reaction with the water. Part of the slowness of the calcium-water reaction results from the metal being partly protected by insoluble white calcium hydroxide. In water solutions of acids where the salt is water soluble, calcium reacts vigorously. Calcium, with a specific mass of 1.55 g/cm3, is the lightest of the alkali earth metals; magnesium is heavier (1.74) and beryllium even more heavy (1.84) despite these two elements being lighter in atomic mass. From strontium on the alkali earth metals get heavier along with the atomic mass. Calcium has a higher resistivity than copper or aluminium. Yet, weight for weight, allowing for its much lower density, it is a rather better conductor than either. However, its use in terrestrial applications is usually limited by its high reactivity with air. Calcium salts are colorless from any contribution of the calcium, and ionic solutions of calcium (Ca2+) are colorless as well. Many calcium salts are not soluble in water. When in solution, the calcium ion to the human taste varies remarkably, being reported as mildly salty, sour, "mineral like" or even "soothing." It is apparent that many animals can taste, or develop a taste, for calcium, and use this sense to detect the mineral in salt licks or other sources.[3] In human nutrition, soluble calcium salts may be added to tart juices without much effect to the average palate. Calcium is the fifth most abundant element by mass in the human body, where it is a common cellular ionic messenger with many functions, and serves also as a structural element in bone. It is the relatively high atomic-numbered calcium in the skeleton which causes bone to be radio-opaque. Of the human body's solid components after drying (as for example, after cremation), about a third of the total mass is the approximately one kilogram of calcium which composes the average skeleton (the remainder being mostly phosphorus and oxygen).
3
Calcium
Occurrence Calcium is not naturally found in its elemental state. Calcium occurs most commonly in sedimentary rocks in the minerals calcite, dolomite and gypsum. It also occurs in igneous and metamorphic rocks chiefly in the silicate minerals: plagioclase, amphiboles, pyroxenes and garnets. See also Calcium minerals.
Applications Some uses are: • as a reducing agent in the extraction of other metals, such as uranium, zirconium, and thorium. • as a deoxidizer, desulfurizer, or decarbonizer for various ferrous and nonferrous alloys. • as an alloying agent used in the production of aluminium, beryllium, copper, lead, and magnesium alloys. • in the making of cements and mortars to be used in construction. • in the making of cheese, where calcium ions influence the activity of rennin in bringing about the coagulation of milk.
Calcium compounds • Calcium carbonate (CaCO3) used in manufacturing cement and mortar, lime, limestone (usually used in the steel industry); aids in production in the glass industry, also has chemical and optical uses as mineral specimens in toothpastes, for example. • Calcium hydroxide solution (Ca(OH)2) (also known as limewater) is used to detect the presence of carbon dioxide by being bubbled through a solution. It turns cloudy where CO2 is present. • Calcium arsenate (Ca3(AsO4)2) is used in insecticides. • Calcium carbide (CaC2) is used: to make acetylene gas (for use in acetylene torches for welding) and in the manufacturing of plastics. • Calcium chloride (CaCl2) is used: in ice removal and dust control on dirt roads, in conditioner for concrete, as an additive in canned tomatoes, and to provide body for automobile tires. • Calcium cyclamate (Ca(C6H11NHSO3)2) was used as a sweetening agent but is no longer permitted for use because of suspected cancer-causing properties. • Calcium gluconate (Ca(C6H11O7)2) is used as a food additive and in vitamin pills. • Calcium hypochlorite (Ca(OCl)2) is used: as a swimming pool disinfectant, as a bleaching agent, as an ingredient in deodorant, and in algaecide and fungicide. • Calcium permanganate (Ca(MnO4)2) is used in liquid rocket propellant, textile production, as a water sterilizing agent and in dental procedures. • Calcium phosphate (Ca3(PO4)2) is used as a supplement for animal feed, fertilizer, in commercial production for dough and yeast products, in the manufacture of glass, and in dental products. • Calcium phosphide (Ca3P2) is used in fireworks, rodenticide, torpedoes and flares. • Calcium stearate (Ca(C18H35O2)2) is used in the manufacture of wax crayons, cements, certain kinds of plastics and cosmetics, as a food additive, in the production of water resistant materials and in the production of paints.
4
Calcium • Calcium sulfate (CaSO4·2H2O) is used as common blackboard chalk, as well as, in its hemihydrate form being more well known as Plaster of Paris. • Calcium tungstate (CaWO4) is used in luminous paints, fluorescent lights and in X-ray studies. • Hydroxylapatite (Ca5(PO4)3(OH), but is usually written Ca10(PO4)6(OH)2) makes up seventy percent of bone. Also carbonated-calcium deficient hydroxylapatite is the main mineral of which dental enamel and dentin are comprised.
H and K lines In the visible portion of the spectrum of many stars, including the Sun, strong absorption lines of singly-ionized calcium are shown. Prominent among these are the H-line at 3968.5 Å and the K line at 3933.7 Å of singly-ionized calcium, or Ca II. For the Sun and stars with low temperatures, the prominence of the H and K lines can be an indication of strong magnetic activity in the chromosphere. Measurement of periodic variations of these active regions can also be used to deduce the rotation periods of these stars.[4]
History Calcium (Latin word calcis meaning "lime") was known as early as the first century when the Ancient Romans prepared lime as calcium oxide. Literature dating back to 975 AD notes that plaster of paris (calcium sulphate), is useful for setting broken bones. It was not isolated until 1808 in England when Sir Humphry Davy electrolyzed a mixture of lime and mercuric oxide. Davy was trying to isolate calcium; when he heard that Swedish chemist Jöns Jakob Berzelius and Pontin prepared calcium amalgam by electrolyzing lime in mercury, he tried it himself. He worked with electrolysis throughout his life and also discovered/isolated sodium, potassium, magnesium, boron and barium. Calcium metal was not available in large scale until the beginning of the 20th century.
Compounds Calcium, combined with phosphate to form hydroxylapatite, is the mineral portion of human and animal bones and teeth. The mineral portion of some corals can also be transformed into hydroxylapatite. Calcium hydroxide (slaked lime) is used in many chemical refinery processes and is made by heating limestone at high temperature (above 825°C) and then carefully adding water to it. When lime is mixed with sand, it hardens into a mortar and is turned into plaster by carbon dioxide uptake. Mixed with other compounds, lime forms an important part of Portland cement. Calcium carbonate (CaCO3) is one of the common compounds of calcium. It is heated to form quicklime (CaO), which is then added to water (H2O). This forms another material known as slaked lime (Ca(OH)2), which is an inexpensive base material used throughout the chemical industry. Chalk, marble, and limestone are all forms of calcium carbonate. When water percolates through limestone or other soluble carbonate rocks, it partially dissolves the rock and causes cave formation and characteristic stalactites and stalagmites and also forms hard water. Other important calcium compounds are calcium nitrate, calcium sulfide, calcium chloride, calcium carbide, calcium cyanamide and calcium hypochlorite.
5
Calcium
6
Isotopes Calcium has four stable isotopes (40Ca and 42Ca through 44Ca), plus two more isotopes (46Ca and 48Ca) that have such long half-lives that for all practical purposes they can be considered stable. It also has a cosmogenic isotope, radioactive 41Ca, which has a half-life of 103,000 years. Unlike cosmogenic isotopes that are produced in the atmosphere, 41Ca is produced by neutron activation of 40Ca. Most of its production is in the upper metre or so of the soil column, where the cosmogenic neutron flux is still sufficiently strong. 41Ca has received much attention in stellar studies because it decays to 41K, a critical indicator of solar-system anomalies. 97% of naturally occurring calcium is in the form of 40Ca. 40Ca is one of the daughter products of 40K decay, along with 40Ar. While K-Ar dating has been used extensively in the geological sciences, the prevalence of 40Ca in nature has impeded its use in dating. Techniques using mass spectrometry and a double spike isotope dilution have been used for K-Ca age dating. The most abundant isotope, 40Ca, has a nucleus of 20 protons and 20 neutrons. This is the heaviest stable isotope of any element which has equal numbers of protons and neutrons. In supernova explosions, calcium is formed from the reaction of carbon with various numbers of alpha particles (helium nuclei), until the most common calcium isotope (containing 10 helium nuclei) has been synthesized.
Nutrition Recommended Adequate Intake by the IOM for Calcium: Age
Calcium (mg/day)
0–6 months
210
7–12 months
270
1–3 years
500
4–8 years
800
9–18 years
1300
19–50 years
1000
51+ years
1200
Calcium is an important component of a healthy diet and a mineral necessary for life. The National Osteoporosis Foundation says, "Calcium plays an important role in building stronger, denser bones early in life and keeping bones strong and healthy later in life." Approximately ninety-nine percent of the body's calcium is stored in the bones and teeth.[5] The rest of the calcium in the body has other important uses, such as some exocytosis, especially neurotransmitter release, and muscle contraction. In the electrical conduction system of the heart, calcium replaces sodium as the mineral that depolarizes the cell, proliferating the action potential. In cardiac muscle, sodium influx commences an action potential, but during potassium efflux, the cardiac myocyte experiences calcium influx, prolonging the action potential and creating a plateau phase of dynamic equilibrium. Long-term calcium deficiency can lead to rickets and poor blood clotting and in case of a menopausal woman, it can lead to osteoporosis, in which the bone deteriorates and there is an increased risk of fractures. While a lifelong deficit can affect bone and tooth formation,
Calcium
7
over-retention can cause hypercalcemia (elevated levels of calcium in the blood), impaired kidney function and decreased absorption of other minerals.[6] High calcium intakes or high calcium absorption were previously thought to contribute to the development of kidney stones. However, a high calcium intake has been associated with a lower risk for kidney [7] [8] [9] stones in more recent research. Vitamin D is needed to absorb calcium. Dairy products, such as milk and cheese, are a well-known source of calcium. However, some individuals are allergic to dairy products and even more people, particularly those of non Indo-European descent, are lactose-intolerant, leaving them unable to consume non-fermented dairy products in quantities larger than about half a liter per serving. Others, such as vegans, avoid dairy products for ethical and health reasons. Fortunately, many good sources of calcium exist. These include seaweeds such as kelp, wakame and hijiki; nuts and seeds (like almonds and sesame); blackstrap molasses; beans; oranges; figs; quinoa; amaranth; collard greens; okra; rutabaga; broccoli; dandelion leaves; kale; and fortified products such as orange juice and soy milk. (However, calcium fortified orange juice often contains vitamin D3 derived from lanolin, and is thus unacceptable for vegans.[10] ) An overlooked source of calcium is eggshell, which can be ground into a powder and mixed into food or a glass of water.[11] [12] [13] Cultivated vegetables generally have less calcium than wild plants.[14] The calcium content of most foods can be found in the USDA National Nutrient Database.[15]
Dietary calcium supplements Calcium supplements are used to prevent and to treat calcium deficiencies. Most experts recommend that supplements be taken with food and that no more than 600 mg should be taken at a time because the percent of calcium absorbed decreases as the amount of calcium in the supplement increases.[] It is recommended to spread doses throughout the day. Recommended daily calcium intake for adults ranges from 1000 to 1500 mg. It is recommended to take supplements with food to aid in absorption.
500 milligram calcium supplements made from calcium carbonate
Vitamin D is added to some calcium supplements. Proper vitamin D status is important because vitamin D is converted to a hormone in the body which then induces the synthesis of intestinal proteins responsible for calcium absorption.[16] • The absorption of calcium from most food and commonly-used dietary supplements is very similar.[17] This is contrary to what many calcium supplement manufacturers claim in their promotional materials. • Milk is an excellent source of dietary calcium because it has a high concentration of calcium and the calcium in milk is excellently absorbed.[17] • Calcium carbonate is the most common and least expensive calcium supplement. It should be taken with food. It depends on low pH levels for proper absorption in the intestine.[18] Some studies suggests that the absorption of calcium from calcium carbonate is similar to the absorption of calcium from milk.[19] [20] While most people digest calcium carbonate very well, some might develop gastrointestinal discomfort or gas. Taking magnesium with it can help to avoid constipation. Calcium carbonate is 40%
Calcium elemental calcium. 1000 mg will provide 400 mg of calcium. However, supplement labels will usually indicate how much calcium is present in each serving, not how much calcium carbonate is present. • Antacids, such as Tums, frequently contain calcium carbonate, and are a very commonly-used, inexpensive calcium supplement. • Coral Calcium is a salt of calcium derived from fossilized coral reefs. Coral calcium is composed of calcium carbonate and trace minerals. • Calcium citrate can be taken without food and is the supplement of choice for individuals with achlorhydria or who are taking histamine-2 blockers or proton-pump inhibitors.[21] It is more easily digested and absorbed than calcium carbonate if taken on empty stomach and less likely to cause constipation and gas than calcium carbonate. It also has a lower risk of contributing to the formation of kidney stones. Calcium citrate is about 21% elemental calcium. 1000 mg will provide 210 mg of calcium. It is more expensive than calcium carbonate and more of it must be taken to get the same amount of calcium. • Calcium phosphate costs more than calcium carbonate, but less than calcium citrate. It is easily absorbed and is less likely to cause constipation and gas than either. • Calcium lactate has similar absorption as calcium carbonate[22] , but is more expensive. Calcium lactate and calcium gluconate are less concentrated forms of calcium and are not practical oral supplements.[21] • Calcium chelates are synthetic calcium compounds, with calcium bound to an organic molecule, such as malate, aspartate, or fumarate. These forms of calcium may be better absorbed on an empty stomach. However, in general they are absorbed similarly to calcium carbonate and other common calcium supplements when taken with food.[23] The 'chelate' mimics the action that natural food performs by keeping the calcium soluble in the intestine. Thus, on an empty stomach, in some individuals, chelates might theoretically be absorbed better. • Microcrystalline hydroxyapatite (MH) is marketed as a calcium supplement, and has in some randomized trials been found to be more effective than calcium carbonate. • Orange juice with calcium added is a good dietary source for persons who have lactose intolerance. In July 2006, a report citing research from Fred Hutchinson Cancer Research Center in Seattle, Washington claimed that women in their 50s gained 5 pounds less in a period of 10 years by taking more than 500 mg of calcium supplements than those who did not. However, the doctor in charge of the study, Dr. Alejandro J. Gonzalez also noted it would be "going out on a limb" to suggest calcium supplements as a weight-limiting aid.[24]
Prevention of fractures due to osteoporosis Such studies often do not test calcium alone, but rather combinations of calcium and vitamin D. Randomized controlled trials found both positive[25] [26] and negative[27] [28] [29] [30] effects. The different results may be explained by doses of calcium and underlying rates of calcium supplementation in the control groups.[31] However, it is clear that increasing the intake of calcium promotes deposition of calcium in the bones, where it is of more benefit in preventing the compression fractures resulting from the osteoporotic thinning of the dendritic web of the bodies of the vertebrae, than it is at preventing the more serious cortical bone fractures which happen at hip and wrist.
8
Calcium
9
Possible cancer prevention A meta-analysis[26] by the international Cochrane Collaboration of two randomized controlled trials[32] [33] found that calcium "might contribute to a moderate degree to the prevention of adenomatous colonic polyps". More recent studies were conflicting, and one which was positive for effect (Lappe, et al.) did control for a possible anti-carcinogenic effect of vitamin D, which was found to be an independent positive influence from calcium-alone on cancer risk (see second study below) [34] . • A randomized controlled trial found that 1000 mg of elemental calcium and 400 IU of vitamin D3 had no effect on colorectal cancer[35] • A randomized controlled trial found that 1400–1500 mg supplemental calcium and 1100 IU vitamin D3 reduced aggregated cancers with a relative risk of 0.402.[36] • An observational cohort study found that high calcium and vitamin D intake was associated with "lower risk of developing premenopausal breast cancer."[37]
Overdose Exceeding the recommended daily calcium intake for an extended period of time can result in hypercalcemia and calcium metabolism disorder.
See also • • • •
Calcium metabolism Calcium in biology Calcium compounds Disorders of calcium metabolism
References • Rebecca J. Donatelle. Health, The Basics. 6th ed. San Francisco: Pearson Education, Inc. 2005.
External links • WebElements.com — Calcium [38] • USDA National Nutrient Database, Calcium content of selected foods • UK Food Standards Agency: Calcium [40]
[39]
Calcium
References [1] A. G. Dickson, C. Goyet (1994). " 5 (http:/ / cdiac. esd. ornl. gov/ ftp/ cdiac74/ chapter5. pdf)". Handbook of method for the analysis of the various parameters of the carbon dioxide system in sea water, version 2. ORNL/CDIAC-74. . [2] Pauling, Linus (1970). General Chemistry. Dover Publications. p. 627. ISBN 0716701499. [3] M. G. Tordoff. " Calcium: Taste, Intake, and Appetite (http:/ / physrev. physiology. org/ cgi/ content/ full/ 81/ 4/ 1567)". Physiological Reviews 81 (4): 1567. . [4] Staff (1995). " H-K Project (http:/ / www. mtwilson. edu/ hk/ )". Mount Wilson Observatory. . Retrieved 2006-08-10. [5] " Osteoporosis Prevention - Calcium Recommendations (http:/ / www. nof. org/ prevention/ calcium2. htm)". . [6] Standing Committee on the Scientific Evaluation of Dietary Reference Intakes, Food and Nutrition Board, Institute of Medicine (1997). Dietary Reference Intakes for Calcium, Phosphorus, Magnesium, Vitamin D and fluoride. Washington DC: The National Academies Press. ISBN 0309064031. [7] Curhan, Gc; Willett, Wc; Rimm, Eb; Stampfer, Mj (Mar 1993). "A prospective study of dietary calcium and other nutrients and the risk of symptomatic kidney stones.". The New England journal of medicine 328 (12): 833–8. ISSN 0028-4793 (http:/ / worldcat. org/ issn/ 0028-4793). PMID 8441427. [8] Bihl G, Meyers A. (August 2001). "Recurrent renal stone disease-advances in pathogenesis and clinical management". Lancet 358 (9282): 651–656. doi: 10.1016/S0140-6736(01)05782-8 (http:/ / dx. doi. org/ 10. 1016/ S0140-6736(01)05782-8). PMID 11530173. [9] Hall WD, Pettinger M, Oberman A, et al. (July 2001). "Risk factors for kidney stones in older women in the Southern United States". Am J Med Sci 322 (1): 12–18. doi: 10.1097/00000441-200107000-00003 (http:/ / dx. doi. org/ 10. 1097/ 00000441-200107000-00003). PMID 11465241. [10] " Sources of vitamin D in orange juice (http:/ / findarticles. com/ p/ articles/ mi_m0FDE/ is_3_23/ ai_n6138556)". . [11] Anne Schaafsma, Gerard M Beelen (1999). " Eggshell powder, a comparable or better source of calcium than purified calcium carbonate: piglet studies (http:/ / www3. interscience. wiley. com/ cgi-bin/ abstract/ 63003036/ ABSTRACT)" (abstract). Journal of the Science of Food and Agriculture 79 (12): 1596–1600. doi: 10.1002/(SICI)1097-0010(199909)79:12<1596::AID-JSFA406>3.0.CO;2-A (http:/ / dx. doi. org/ 10. 1002/ (SICI)1097-0010(199909)79:12<1596::AID-JSFA406>3. 0. CO;2-A). . [12] Schaafsma A, van Doormaal JJ, Muskiet FA, Hofstede GJ, Pakan I, van der Veer E (March 2002). "Positive effects of a chicken eggshell powder-enriched vitamin-mineral supplement on femoral neck bone mineral density in healthy late post-menopausal Dutch women". Br. J. Nutr. 87 (3): 267–75. doi: 10.1079/BJNBJN2001515 (http:/ / dx. doi. org/ 10. 1079/ BJNBJN2001515). PMID 12064336. [13] Rovenský J, Stancíková M, Masaryk P, Svík K, Istok R (2003). "Eggshell calcium in the prevention and treatment of osteoporosis". Int J Clin Pharmacol Res 23 (2-3): 83–92. PMID 15018022. [14] " Original Wild Foods vs. Available Foods Today for Instinctos (http:/ / www. beyondveg. com/ nieft-k/ instincto-guide/ instincto-guide1e. shtml)". . [15] " USDA National Nutrient Database (http:/ / www. nal. usda. gov/ fnic/ foodcomp/ search)". . [16] Combs, G (2008). The Vitamins. Academic Press. p. 161. [17] Weaver, CM (2006). "Calcium". Present Knowledge in Nutrition, 9th Ed.. I. ILSI Press. p. 377. [18] Remington, Joseph (2005). Remington: The Science and Practice of Pharmacy. Lippincott Williams & Wilkins. pp. 1338. ISBN 0781746736. [19] Zhao, Y. et al.. "Calcium bioavailability of calcium carbonate fortified soy milk is equivalent to cow's milk in young women". J. Nutr. 135 (10): 2379. [20] Ligia Martini and Richard J Wood (2002). " Relative bioavailability of calcium-rich dietary sources in the elderly (http:/ / www. ajcn. org/ cgi/ content/ abstract/ 76/ 6/ 1345)". American Journal of Clinical Nutrition 76 (6): 1345–1350. . [21] Straub, D. A. (2007). "Calcium Supplementation in Clinical Practice: A Review of Forms, Doses, and Indications". Nutrition in Clinical Practice 22: 286. doi: 10.1177/0115426507022003286 (http:/ / dx. doi. org/ 10. 1177/ 0115426507022003286). [22] Martin, Berdine R. (2002). "Calcium Absorption from Three Salts and CaSO4-Fortified Bread in Premenopausal Women". Journal of Agricultural and Food Chemistry 50: 3874. doi: 10.1021/jf020065g (http:/ / dx. doi. org/ 10. 1021/ jf020065g). [23] Weaver, Connie M. (2002). "Absorption of Calcium Fumarate Salts Is Equivalent to Other Calcium Salts When Measured in the Rat Model". Journal of Agricultural and Food Chemistry 50: 4974. doi: 10.1021/jf0200422 (http:/ / dx. doi. org/ 10. 1021/ jf0200422). [24] Anne Harding. " Calcium May Help With Weight Loss (http:/ / www. rxalternativemedicine. com/ headlines_news. php#headline77)". . Retrieved 2007-07-10.
10
Calcium [25] Dawson-Hughes B, Harris SS, Krall EA, Dallal GE (1997). "Effect of calcium and vitamin D supplementation on bone density in men and women 65 years of age or older". N. Engl. J. Med. 337 (10): 670–6. doi: 10.1056/NEJM199709043371003 (http:/ / dx. doi. org/ 10. 1056/ NEJM199709043371003). PMID 9278463. [26] Weingarten MA, Zalmanovici A, Yaphe J (2005). "Dietary calcium supplementation for preventing colorectal cancer, adenomatous polyps and calcium metabolisism disorder.". Cochrane database of systematic reviews (Online) (3): CD003548. doi: 10.1002/14651858.CD003548.pub3 (http:/ / dx. doi. org/ 10. 1002/ 14651858. CD003548. pub3). PMID 16034903. [27] Jackson RD, LaCroix AZ, Gass M, et al. (2006). "Calcium plus vitamin D supplementation and the risk of fractures". N. Engl. J. Med. 354 (7): 669–83. doi: 10.1056/NEJMoa055218 (http:/ / dx. doi. org/ 10. 1056/ NEJMoa055218). PMID 16481635. [28] Grant AM, Avenell A, Campbell MK, et al. (2005). "Oral vitamin D3 and calcium for secondary prevention of low-trauma fractures in elderly people (Randomised Evaluation of Calcium Or vitamin D, RECORD): a randomised placebo-controlled trial". Lancet 365 (9471): 1621–8. doi: 10.1016/S0140-6736(05)63013-9 (http:/ / dx. doi. org/ 10. 1016/ S0140-6736(05)63013-9). PMID 15885294. [29] Porthouse J, Cockayne S, King C, et al. (2005). "Randomised controlled trial of calcium and supplementation with cholecalciferol (vitamin D3) for prevention of fractures in primary care". BMJ 330 (7498): 1003. doi: 10.1136/bmj.330.7498.1003 (http:/ / dx. doi. org/ 10. 1136/ bmj. 330. 7498. 1003). PMID 15860827. [30] Prince RL, Devine A, Dhaliwal SS, Dick IM (2006). "Effects of calcium supplementation on clinical fracture and bone structure: results of a 5-year, double-blind, placebo-controlled trial in elderly women". Arch. Intern. Med. 166 (8): 869–75. doi: 10.1001/archinte.166.8.869 (http:/ / dx. doi. org/ 10. 1001/ archinte. 166. 8. 869). PMID 16636212. [31] Fletcher RH (2006). " Calcium plus vitamin D did not prevent hip fracture or colorectal cancer in postmenopausal women (http:/ / www. acpjc. org/ Content/ 145/ 1/ issue/ ACPJC-2006-145-1-004. htm)" (subscription required). ACP J. Club 145 (1): 4–5. PMID 16813354. . [32] Baron JA, Beach M, Mandel JS, et al. (1999). "Calcium supplements for the prevention of colorectal adenomas. Calcium Polyp Prevention Study Group". N. Engl. J. Med. 340 (2): 101–7. doi: 10.1056/NEJM199901143400204 (http:/ / dx. doi. org/ 10. 1056/ NEJM199901143400204). PMID 9887161. [33] Bonithon-Kopp C, Kronborg O, Giacosa A, Räth U, Faivre J (2000). "Calcium and fibre supplementation in prevention of colorectal adenoma recurrence: a randomised intervention trial. European Cancer Prevention Organisation Study Group". Lancet 356 (9238): 1300–6. doi: 10.1016/S0140-6736(00)02813-0 (http:/ / dx. doi. org/ 10. 1016/ S0140-6736(00)02813-0). PMID 11073017. [34] Lappe, Jm; Travers-Gustafson, D; Davies, Km; Recker, Rr; Heaney, Rp (Jun 2007). " Vitamin D and calcium supplementation reduces cancer risk: results of a randomized trial. (http:/ / www. ajcn. org/ cgi/ pmidlookup?view=long& pmid=17556697)" (Free full text). The American journal of clinical nutrition 85 (6): 1586–91. ISSN 0002-9165 (http:/ / worldcat. org/ issn/ 0002-9165). PMID 17556697. . [35] Wactawski-Wende J, Kotchen JM, Anderson GL, et al. (2006). "Calcium plus vitamin D supplementation and the risk of colorectal cancer". N. Engl. J. Med. 354 (7): 684–96. doi: 10.1056/NEJMoa055222 (http:/ / dx. doi. org/ 10. 1056/ NEJMoa055222). PMID 16481636. [36] Lappe JM, Travers-Gustafson D, Davies KM, Recker RR, Heaney RP (2007). "Vitamin D and calcium supplementation reduces cancer risk: results of a randomized trial". Am. J. Clin. Nutr. 85 (6): 1586–91. PMID 17556697. [37] Lin J, Manson JE, Lee IM, Cook NR, Buring JE, Zhang SM (2007). "Intakes of calcium and vitamin d and breast cancer risk in women". Arch. Intern. Med. 167 (10): 1050–9. doi: 10.1001/archinte.167.10.1050 (http:/ / dx. doi. org/ 10. 1001/ archinte. 167. 10. 1050). PMID 17533208. [38] http:/ / www. webelements. com/ webelements/ elements/ text/ Ca/ index. html [39] http:/ / www. nal. usda. gov/ fnic/ foodcomp/ Data/ SR17/ wtrank/ sr17a301. pdf [40] http:/ / www. eatwell. gov. uk/ healthydiet/ nutritionessentials/ vitaminsandminerals/ calcium/
11
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12
Scandium
1
Scandium calcium ← scandium → titanium ↑ Sc ↓ Y
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
21Sc Periodic table
Appearance silvery white
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
1645
1804
(2006)
(2266)
(2613)
(3101)
Scandium
2
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1235.0 kJ·mol−1 3rd: 2388.6 kJ·mol−1Atomic radiusCovalent radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityThermal expansionYoung's modulusShear modulusBulk modulusPoisson ratioBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of scandium iso 44m
Sc
N.A. syn
half-life 58.61 h
45
100%
45
46
syn
83.79 d
Sc Sc
47
Sc
48
Sc
syn
syn
DM
DE (MeV)
DP
IT
0.2709
44
γ
1.0, 1.1, 1.1
44
ε
-
44
β−
0.3569
46
γ
0.889, 1.120
-
β−
0.44, 0.60
47
γ
0.159
-
β−
0.661
48
γ
0.9, 1.3, 1.0
-
Sc Sc Ca
Sc is stable with 24 neutron
3.3492 d
43.67 h
Ti
Ti
Ti
scandium, Sc, 21 transition metal3, 4, d44.955912(6) g·mol−1 [Ar] 3d1 4s2 2, 8, 9, 2 (Image) solid 2.985 g·cm−3 2.80 g·cm−3 1814 K,1541 °C,2806 °F 3109 K,2836 °C,5136 °F 14.1 kJ·mol−1 332.7 kJ·mol−1 (25 °C) 25.52 J·mol−1·K−13, 2[1] , 1 [2] (weakly basic oxide) 1.36 (Pauling scale) 1st: 633.1 kJ·mol−1162 pm170±7 pm hexagonal paramagnetic (r.t.) (α, poly) calc. 562 nΩ·m (300 K) 15.8 W·m−1·K−1 (r.t.) (α, poly) 10.2 µm/(m·K) 74.4 GPa 29.1 GPa 56.6 GPa 0.279 750 MPa 7440-20-2 Scandium (pronounced /ˈskændiəm/) is a chemical element with symbol Sc and atomic number 21. A silvery-white metallic transition metal, it has historically been sometimes classified as a rare earth element, together with yttrium and the lanthanides. In 1879 Lars Fredrik Nilson and his team, found a new element with spectral analysis, in the minerals euxenite and gadolinite from Scandinavia. Scandium is present in most of the rare earth element and uranium deposits, but it is extracted from these ores in only a few mines worldwide. Due to the low availability and the difficulties in the preparation of metallic scandium, which was first done in 1937, it took until the 1970s before applications for scandium were developed. The positive effects of scandium on aluminium alloys were discovered in the 1970s, and its use in such alloys remains the only major application of scandium.
Scandium
3
History Dmitri Mendeleev predicted the existence of an element that he called ekaboron, with an atomic mass between 40 and 48 in 1869. Ten years later Lars Fredrik Nilson found a new element in the minerals euxenite and gadolinite from Scandinavia. He was able to prepare 2 g of scandium oxide of high purity. [3] [4] He named it scandium, from the Latin Scandia meaning "Scandinavia". Nilson was apparently unaware of Mendeleev's prediction, but Per Teodor Cleve recognized the correspondence and notified Mendeleev.[5] Metallic scandium was produced for the first time in 1937 by electrolysis of a eutectic mixture, at 700–800 °C, of potassium, lithium, and scandium chlorides.[6] The first pound of 99% pure scandium metal was produced in 1960. The use for aluminium alloys began in 1971, following a US patent. Aluminium-scandium alloys were also developed in the USSR.[7]
Position in the periodic table Groups 1 to 3 of the periodic table could be written as follows: 1
2
3
H Li
Be
B
Na
Mg
Al
K
Ca
Sc
Rb
Sr
Y
Cs
Ba
La
Fr
Ra
Ac
This grouping is consistent with Mendeleev's prediction for scandium as "eka-boron". It shows that the properties of Sc will be intermediate between the properties of Al and Y, in the same way that the properties of Ca are intermediate between those of Mg and Sr. It also shows that there will be a diagonal relationship between Mg and Sc, just as there is between Be and Al. However, in the standard periodic table boron and aluminium are placed in group 13, where the relationships above are less obvious. As to the rest of group 3, there has been controversy as to whether yttrium is in the same group as lanthanum or as lutetium.[8] In the chemical compounds of the elements shown as group 3, above, the predominant oxidation state is +3. The ions M3+ will all have the electronic configuration of a noble gas, so it is reasonable that they should be in the same group of the periodic table. Most modern text-books place Sc, Y, La and Ac in the same periodic group.
Occurrence Scandium does not have a particularly low abundance in the earth's crust. Estimates vary from 18 to 25 ppm, which is comparable to the abundance of cobalt (20–30 ppm). However, scandium is distributed sparsely and occurs in trace amounts in many minerals.[9] Rare minerals from Scandinavia[10] and Madagascar[11] such as thortveitite, euxenite, and gadolinite are the only known concentrated sources of this element. Thortveitite can contain up to 45%, as scandium(III) oxide.[10]
Scandium Scandium is more common in the sun and certain stars than on Earth. Scandium is only the 50th most common element on earth (35th most abundant in the Earth's crust), but it is the 23rd most common element in the sun.[12]
Production World production of scandium is in the order of 2,000 kg per year as scandium oxide. The primary production is 400 kg while the rest is from stockpiles of Russia created during the Cold War. In 2003 only three mines produced scandium: the uranium and iron mines in Zhovti Vody in Ukraine, the rare earth mines in Bayan Obo, China and the apatite mines in the Kola peninsula, Russia. In each case scandium is a byproduct from the extraction of other elements.[13] and is sold as scandium oxide. The production of metallic scandium is in the order of 10 kg per year.[13] [14] The oxide is converted to scandium fluoride and reduced with metallic calcium. Madagascar and Iveland-Evje Region in Norway have the only deposits of minerals with high scandium content, thortveitite (Y,Sc)2(Si2O7) and kolbeckite ScPO4·2H2O, but these are not being exploited.[14] Other scandium sources include the nickel and cobalt mines at Syerston and Lake Innes, New South Wales, Australia, iron, tin, and tungsten deposits in China and uranium deposits in Russia and Kazakhstan. As of 2003, scandium was not being extracted from the tailings at any of these mines, but some scandium extraction may be started if there is sufficient demand.[13] There is currently no primary production of scandium in the Americas, Europe, or Australia.
Isotopes Naturally occurring scandium is composed of one stable isotope 45Sc with a nuclear spin of 7/2. 13 radioisotopes have been characterized with the most stable being 46Sc with a half-life of 83.8 days, 47Sc with a half-life of 3.35 days, and 48Sc with a half-life of 43.7 hours. All of the remaining radioactive isotopes have half lives that are less than 4 hours, and the majority of these have half-lives that are less than 2 minutes. This element also has 5 meta states with the most stable being 44mSc (t½ 58.6 h).[15] The isotopes of scandium range in atomic weight from 40 u (40Sc) to 54 u (54Sc). The primary decay mode at masses lower than the only stable isotope, 45Sc, is electron capture, and the primary mode at masses above it is beta emission. The primary decay products at atomic weights below 45Sc are calcium isotopes and the primary products from higher atomic weights are titanium isotopes.[15]
Compounds Scandium metal is hard and has a silvery appearance. It develops a slightly yellowish or pinkish cast when exposed to air. It is not resistant to weathering and dissolves slowly in most dilute acids. It does not react with a 1:1 mixture of nitric acid (HNO3) and hydrofluoric acid, HF, presumably due to the formation of an impermeable passive layer on the surface of the metal. In the compounds ScB and ScC, boron and carbon are incorporated non-stoichiometrically into the lattice of the scandium.[16] The radii of M3+ ions in the following table
4
Scandium
5
Ionic radii (pm) Al
Sc
Y
La
Lu
53.5
74.5
90.0
103.2
86.1
indicate why the chemistry of scandium is more closely related to that of yttrium than that of aluminium and explains why scandium has been classified as a lanthanide-like element. The oxide Sc2O3 is weakly acidic and the hydroxide Sc(OH)3 is amphoteric Sc3+ (aq.) ← H+ + Sc(OH)3 + OH− → Sc(OH)
The α- and γ - forms of scandium oxide hydroxide (ScO(OH)), are isostructural with their aluminium oxide hydroxide counterparts.[17] Solutions of Sc3+ in water are acidic because of hydrolysis. The halides ScX3 (X = Cl, Br, I) are very soluble in water, but ScF3 is insoluble. In all four halides the scandium is 6-coordinate. The halides are Lewis acids; for example, ScF3 dissolves a solution containing excess fluoride to form [ScF6]3−. This is a typical example of a complex of Sc(III) in which the coordination number is 6. In the larger Y and La ions 8and 9- coordination are often found. There are a few compounds known in which the oxidation state is less than 3. The cluster [Sc6Cl12]3− has a similar structure to that of the Nb6Cl12 cluster in which chlorine atoms bridge the 12 edges of an octahedron of metal atoms.[18] Other sub-halides are known. The nature of the hydride ScH2 is not yet fully understood.[2] It appears not to be a saline hydride of Sc(II), but may be a compound of Sc(III) with two hydrides and an electron which is delocalized in a kind of metallic structure. ScH can be observed spectroscopically at high temperatures in the gas phase.[1] Scandium forms a series of organometallic compounds with cyclopentadienyl, based on the Sc(Cp)2 motif. The chlorine-bridged dimer, [Sc(Cp)2Cl]2 is the starting point for the preparation of many compounds by replacement of the chlorine.[19]
Applications The addition of scandium to aluminium limits the excessive grain growth that occurs in the heat-affected zone of welded aluminium components. This has two beneficial effects: the precipitated Al3Sc forms smaller crystals than are formed in other aluminium alloys[20] and the volume of precipitate-free zones that normally exist at the grain boundaries of age-hardening aluminium alloys is reduced.[21] Both of these effects increase the usefulness of the alloy. However, titanium alloys, which are similar in lightness and strength, are cheaper and much more widely used.[22]
Parts of the Mig–29 are made from [20] Al-Sc alloy.
The main application of scandium by weight is in aluminium-scandium alloys for minor aerospace industry components. These alloys contain between 0.1% and 0.5% of scandium. They were used in the Russian military aircraft Mig 21 and Mig 29.[21] Some items of sports equipment, which rely on high performance materials, have been made with scandium-aluminium alloys, including baseball bats[23] , lacrosse sticks, as well
Scandium
6
as bicycle[24] frames and components. U.S. gunmaker Smith & Wesson produces revolvers with frames composed of scandium alloy and cylinders of titanium .[25] Approximately 20 kg (as Sc2O3) of scandium is used annually in the United States to make high-intensity discharge lamps.[26] Scandium iodide, along with Sodium Iodide, when added to a modified form of mercury-vapor lamp, produces a form of metal halide lamp, an artificial light source which produce a very white light with high colour rendering index that sufficiently resembles sunlight to allow good color-reproduction with TV cameras.[27] About 80 kg of scandium is used in metal halide lamps/light bulbs globally per year. The first scandium based metal halide lamps were patented by General Electric and initially made in North America, although they are now produced in all major industrialized countries. The radioactive isotope 46Sc is used in oil refineries as a tracing agent.[26] Scandium triflate is a catalytic Lewis acid used in organic chemistry.[28]
Health and safety The pure metal is not considered to be toxic. Little animal testing of scandium compounds has been done.[29] The median lethal dose (LD50) levels for scandium(III) chloride for rats have been determined and were intraperitoneal 4 mg/kg and oral 755 mg/kg.[30] In the light of these results compounds of scandium should be handled as compounds of moderate toxicity.
See also • • • •
Scandium compounds Scandium minerals Yttrium Rare earth element
External links • WebElements.com – Scandium
[31]
References [1] McGuire, Joseph C.; Kempter, Charles P. (1960). "Preparation and Properties of Scandium Dihydride". Journal of Chemical Physics: 1584–1585. doi: 10.1063/1.1731452 (http:/ / dx. doi. org/ 10. 1063/ 1. 1731452). [2] Smith, R. E. (1973). "Diatomic Hydride and Deuteride Spectra of the Second Row Transition Metals". Proceedings of the Royal Society of London. Series A, Mathematical and Physical Sciences (1588): 113–127. doi: 10.1098/rspa.1973.0015 (http:/ / dx. doi. org/ 10. 1098/ rspa. 1973. 0015). [3] Lars Fredrik Nilson (1879). " Sur l'ytterbine, terre nouvelle de M. Marignac (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k30457/ f639. table)". Comptes Rendus 88: 642–647. . [4] F. L. Nilson (1879). "Ueber Scandium, ein neues Erdmetall". Berichte der deutschen chemischen Gesellschaft 12 (1): 554–557. doi: 10.1002/cber.187901201157 (http:/ / dx. doi. org/ 10. 1002/ cber. 187901201157). [5] Per Teodor Cleve (1879). " Sur le scandium (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k3046j/ f432. table)". Comptes Rendus 89: 419–422. . [6] Fischer, Werner; Brünger, Karl; Grieneisen, Hans (1937). "Über das metallische Scandium". Zeitschrift für anorganische und allgemeine Chemie 231 (1-2): 54–62. doi: 10.1002/zaac.19372310107 (http:/ / dx. doi. org/ 10. 1002/ zaac. 19372310107). [7] Zakharov, V. V. (2003). "Effect of Scandium on the Structure and Properties of Aluminum Alloys". Metal Science and Heat Treatment 45: 246. doi: 10.1023/A:1027368032062 (http:/ / dx. doi. org/ 10. 1023/ A:1027368032062). [8] Lavelle, L. (2008). "Lanthanum (La) and Actinium (Ac) Should Remain in the d-block". J. Chem. Ed. 85: 1482.
Scandium [9] Bernhard, F. (2001). "Scandium mineralization associated with hydrothermal lazurite-quartz veins in the Lower Austroalpie Grobgneis complex, East Alps, Austria". Mineral Deposits in the Beginning of the 21st Century. Lisse: Balkema. ISBN 9026518463. [10] Kristiansen, Roy (2003). " Scandium - Mineraler I Norge (http:/ / www. nags. net/ Stein/ 2003/ Sc-minerals. pdf)" (in Norwegian). Stein: 14–23. . [11] von Knorring, O.; Condliffe, E. (1987). "Mineralized pegmatites in Africa". Geological Journal 22: 253. doi: 10.1002/gj.3350220619 (http:/ / dx. doi. org/ 10. 1002/ gj. 3350220619). [12] Lide, David R. (2004). CRC Handbook of Chemistry and Physics. Boca Raton: CRC Press. pp. 4–28. ISBN 9780849304859. [13] Deschamps, Y.. " Scandium (http:/ / www. mineralinfo. org/ Substance/ Scandium/ Sc. pdf)". mineralinfo.com. . Retrieved 2008-10-21. [14] " Mineral Commodity Summaries 2008: Scandium (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ scandium/ mcs-2008-scand. pdf)". United States Geological Survey. . Retrieved 2008-10-20. [15] Audi, Georges (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3–128. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [16] Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils (1985). Lehrbuch der Anorganischen Chemie (91–100 ed.). Walter de Gruyter. pp. 1055–1056. ISBN 3110075113. [17] Christensen, A. Nørlund; Stig Jorgo Jensen (1967). "Hydrothermal Preparation of alpha-ScOOH and of gamma-ScOOH. Crystal Structure of alpha-ScOOH.". Acta Chemica Scandinavica 21: 1121–126.. doi: 10.3891/acta.chem.scand.21-0121 (http:/ / dx. doi. org/ 10. 3891/ acta. chem. scand. 21-0121). [18] Corbett, J.D. (1981). "Extended metal-metal bonding in halides of the early transition metals". Acc. Chem. Res. 14: 239–246. doi: 10.1021/ar00068a003 (http:/ / dx. doi. org/ 10. 1021/ ar00068a003). [19] Marks, T.J.; Ernst, R.D. (1982). "Chapter 21". Comprehensive Organometallic Chemistry (Pergamon Press) 3: 173–270. [20] Ahmad, Zaki (2003). "The properties and application of scandium-reinforced aluminum". JOM 55: 35. doi: 10.1007/s11837-003-0224-6 (http:/ / dx. doi. org/ 10. 1007/ s11837-003-0224-6). [21] Ahmad, Zaki (2003). "The properties and application of scandium-reinforced aluminum". JOM 55: 35. doi: 10.1007/s11837-003-0224-6 (http:/ / dx. doi. org/ 10. 1007/ s11837-003-0224-6). [22] ed. by James A. Schwarz .... (2004). James A. Schwarz, Cristian I. Contescu, Karol Putyera (http:/ / books. google. com/ books?id=aveTxwZm40UC& pg=PA2274). CRC Press. p. 2274. ISBN 0824750497. . [23] Bjerklie, Steve (2006). "A batty business: Anodized metal bats have revolutionized baseball. But are finishers losing the sweet spot?". Metal Finishing 104: 61. doi: 10.1016/S0026-0576(06)80099-1 (http:/ / dx. doi. org/ 10. 1016/ S0026-0576(06)80099-1). [24] " Easton Technology Report : Materials / Scandium (http:/ / www. eastonbike. com/ downloadable_files_unprotected/ r& d_files/ R& D-03 Scandium. pdf)". EastonBike.com. . Retrieved 2009-04-03. [25] " Small Frame (J) - Model 340PD Revolver (http:/ / www. smith-wesson. com/ webapp/ wcs/ stores/ servlet/ ProductDisplay?storeId=10001& catalogId=11101& langId=-1& productId=14765& tabselected=tech& isFirearm=Y& parent_category_rn=15704)". Smith & Wesson. . Retrieved 2008-10-20. [26] C.R. Hammond in CRC Handbook of Chemistry and Physics 85th ed., Section 4; The Elements [27] Simpson, Robert S. (2003). Lighting Control: Technology and Applications (http:/ / books. google. com/ books?id=GEIhCl2T-2EC& pg=PT147& ). Focal Press. pp. 108. ISBN 9780240515663. . [28] Kobayashi, Shu; Manabe, Kei (2000). " Green Lewis acid catalysis in organic synthesis (http:/ / www. iupac. org/ publications/ pac/ 2000/ 7207/ 7207pdf/ 7207kobayashi_1373. pdf)". Pure Appl. Chem. 72 (7): 1373–1380. doi: 10.1351/pac200072071373 (http:/ / dx. doi. org/ 10. 1351/ pac200072071373). . [29] Horovitz, Chaim T.; Birmingham, Scott D. (1999). Biochemistry of Scandium and Yttrium (http:/ / books. google. com/ books?id=1ZTQlCWKjmgC). Springer. ISBN 9780306456572. . [30] Haley, Thomas J.; Komesu, L.; Mavis, N.; Cawthorne, J.; Upham, H. C. (1962). "Pharmacology and toxicology of scandium chloride". Journal of Pharmaceutical Sciences 51: 1043. doi: 10.1002/jps.2600511107 (http:/ / dx. doi. org/ 10. 1002/ jps. 2600511107). [31] http:/ / www. webelements. com/ webelements/ elements/ text/ Sc/ index. html
7
Article Sources and Contributors
Article Sources and Contributors Scandium Source: http://en.wikipedia.org/w/index.php?oldid=305719780 Contributors: 129.132.139.xxx, 16189, 3897515, A Softer Answer, Acroterion, Ahoerstemeier, Aitias, Alansohn, Ale jrb, Andy M. Wang, Anonymous Dissident, Apparition11, Archimerged, AssegaiAli, Aussiepete, Axiosaurus, Baccyak4H, Backslash Forwardslash, Beetstra, Belovedfreak, Benbest, Benjiboi, Bgs022, Bigwhiteyeti, Biochemnick, BlueEarth, Bobo192, Bobthebuilder34, Bomac, Borislav Dopudja, Brian0918, Brockert, Bryan Derksen, C.A.T.S. CEO, CYD, Calvin 1998, Canderson7, CanisRufus, Carnildo, CatherineMunro, Chameleon, Chrisvls, Conversion script, Crystal whacker, DMacks, Dale101usa, Darrien, David Latapie, Deglr6328, Delta G, DocWatson42, Donarreiskoffer, Dwmyers, EH74DK, EddEdmondson, Edgar181, El, El C, Emperorbma, Enigmaman, Eog1916, Epbr123, Evand, Everyguy, Fawcett5, Femto, Fieldday-sunday, Fred Bauder, Gaius Cornelius, Globe Collector, Grendelkhan, Hellbus, Icairns, Ideyal, Ishikawa Minoru, J.delanoy, JNW, Jaan513, Janke, Jaraalbe, Jennavecia, Jerzy, Jj137, Joanjoc, John, John254, Jose77, Junglecat, Jwy, Kajasudhakarababu, Karl-Henner, Kjlewis, Kwamikagami, LA2, LarryMorseDCOhio, Lawlerm, LittleOldMe, MZMcBride, Marcika, MarkV, Maryjosh, Materialscientist, Mav, McSly, Mdf, Mgiganteus1, Mgimpel, Michael Devore, Minesweeper, Mixofall, Montrealais, Mr. Lefty, N2e, Needlenose, Nergaal, Neural, Nick Y., NickW557, Nihiltres, Oxymoron83, Pablothegreat85, Paraballo, Persian Poet Gal, Petergans, Philippe, Piperh, Plexust, Poolkris, Ppanzini, Pras, Psyche825, RTC, Remember, Reyk, Reza kalani, Riana, Rifleman 82, Rolinator, Romanm, RucasHost, SDC, Samuelsen, Saperaud, Satori Son, Schneelocke, Sdsds, Senatorpjt, Sengkang, Serpent's Choice, Shafei, Sjakkalle, Sl, Smallweed, Squids and Chips, Steve Hart, Stewartadcock, Stone, Suisui, Sunborn, Syd Henderson, Tagishsimon, Terra Xin, The Rambling Man, TheNewPhobia, Thefifthamendment, Theseeker4, Thricecube, Tranquility, Vancouverguy, Vsmith, Vuong Ngan Ha, Warrior123w, Warut, Watch37264, Wik, Work permit, Yekrats, Yilloslime, Yumi Kitsuna, Yyy, Zach4636, 364 anonymous edits
Image Sources, Licenses and Contributors file:hexagonal.svg Source: http://en.wikipedia.org/w/index.php?title=File:Hexagonal.svg License: BSD Contributors: Original uploader was Danieljamesscott at en.wikipedia file:Electron shell 021 Scandium.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_021_Scandium.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 file:Scandium.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Scandium.jpg License: GNU Free Documentation License Contributors: Saperaud Image:Mig-29 on landing.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Mig-29_on_landing.jpg License: unknown Contributors: User:Dmottl
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8
Titanium
1
Titanium scandium ← titanium → vanadium ↑ Ti ↓ Zr
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
22Ti Periodic table
Appearance silvery grey-white metallic
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
1982
2171
(2403)
2692
3064
3558
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1309.8 kJ·mol−1 3rd: 2652.5 kJ·mol−1Atomic radiusCovalent radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal
Titanium
2
conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusPoisson ratioMohs hardnessVickers hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of titanium iso 44
Ti
N.A. syn
half-life 63 y
46
8.0%
46
47
7.3%
47
48
73.8%
48
49
5.5%
49
50
5.4%
50
Ti Ti Ti Ti Ti
DM
DE (MeV)
DP
ε
-
44
γ
0.07D, 0.08D
-
Sc
Ti is stable with 24 neutron Ti is stable with 25 neutron Ti is stable with 26 neutron Ti is stable with 27 neutron Ti is stable with 28 neutron
titanium, Ti, 22 transition metal4, 4, d47.867(1) g·mol−1 [Ar] 3d2 4s2 2, 8, 10, 2 (Image) solid 4.506 g·cm−3 4.11 g·cm−3 1941 K,1668 °C,3034 °F 3560 K,3287 °C,5949 °F 14.15 kJ·mol−1 425 kJ·mol−1 (25 °C) 25.060 J·mol−1·K−14, 3, 2, 1[1] (amphoteric oxide) 1.54 (Pauling scale) 1st: 658.8 kJ·mol−1147 pm160±8 pm hexagonal paramagnetic (20 °C) 0.420 µΩ·m (300 K) 21.9 W·m−1·K−1 (25 °C) 8.6 µm·m−1·K−1 (r.t.) 5,090 m·s−1 116 GPa 44 GPa 110 GPa 0.32 6.0 970 MPa 716 MPa 7440-32-6 Titanium (pronounced /taɪˈteɪniəm/) is a chemical element with the symbol Ti and atomic number 22. Sometimes called the “space age metal”, it has a low density and is a strong, lustrous, corrosion-resistant (including to sea water, aqua regia and chlorine) transition metal with a silver color. Titanium can be alloyed with iron, aluminium, vanadium, molybdenum, among other elements, to produce strong lightweight alloys for aerospace (jet engines, missiles, and spacecraft), military, industrial process (chemicals and petro-chemicals, desalination plants, pulp, and paper), automotive, agri-food, medical prostheses, orthopedic implants, dental and endodontic instruments and files, dental implants, sporting goods, jewelry, mobile phones, and other applications.[2] Titanium was discovered in England by William Gregor in 1791 and named by Martin Heinrich Klaproth for the Titans of Greek mythology. The element occurs within a number of mineral deposits, principally rutile and ilmenite, which are widely distributed in the Earth's crust and lithosphere, and it is found in almost all living things, rocks, water bodies, and soils.[2] The metal is extracted from its principal mineral ores via the Kroll process[3] or the Hunter process. Its most common compound, titanium dioxide, is used in the manufacture of white pigments.[4] Other compounds include titanium tetrachloride (TiCl4) (used in smoke screens/skywriting and as a catalyst) and titanium trichloride (TiCl3) (used as a catalyst in the production of polypropylene).[2] The two most useful properties of the metal form are corrosion resistance and the highest strength-to-weight ratio of any metal.[5] In its unalloyed condition, titanium is as strong as some steels, but 45% lighter.[6] There are two allotropic forms[7] and five naturally occurring isotopes of this element; 46Ti through 50Ti, with 48Ti being the most abundant (73.8%).[8] Titanium's properties are chemically and physically similar to zirconium.
Titanium
Characteristics Physical A metallic element, titanium is recognized for its high strength-to-weight ratio.[7] It is a strong metal with low density that is quite ductile (especially in an oxygen-free environment),[9] lustrous, and metallic-white in color.[10] The relatively high melting point (over 1,649 °C or 3,000 °F) makes it useful as a refractory metal. Commercial (99.2% pure) grades of titanium have ultimate tensile strength of about 63,000 psi (434 MPa), equal to that of common, low-grade steel alloys, but are 45% lighter.[6] Titanium is 60% more dense than aluminium, but more than twice as strong[6] as the most commonly used 6061-T6 aluminium alloy. Certain titanium alloys (e.g., Beta C) achieve tensile strengths of over 200000 psi (1400 MPa).[11] However, titanium loses strength when heated above 430 °C (806 °F).[12] It is fairly hard although not as hard as some grades of heat-treated steel, non-magnetic and a poor conductor of heat and electricity. Machining requires precautions, as the material will soften and gall if sharp tools and proper cooling methods are not used. Like those made from steel, titanium structures have a fatigue limit which guarantees longevity in some applications.[10] The metal is a dimorphic allotrope with the hexagonal alpha form changing into the body-centered cubic (lattice) beta form at 882 °C (1620 °F).[12] The specific heat of the alpha form increases dramatically as it is heated to this transition temperature but then falls and remains fairly constant for the beta form regardless of temperature.[12] Similar to zirconium and hafnium, an additional omega phase exists, which is thermodynamically stable at high pressures, but which may exist metastably at ambient pressures. This phase is usually hexagonal (ideal) or trigonal (distorted) and can be viewed as being due to a soft longitudinal acoustic phonon of the beta phase causing collapse of (111) planes of atoms.[13]
Chemical The most noted chemical property of titanium is its excellent resistance to corrosion; it is almost as resistant as platinum, capable of withstanding attack by acids, moist chlorine in water but is soluble in concentrated acids.[14] While the following pourbaix diagram shows that titanium is thermodynamically a very reactive metal, it is slow to react with water and air.
3
Titanium
4 This metal forms a passive and protective oxide coating (leading to increased corrosion-resistance) when exposed to elevated temperatures in air, but at room temperatures it resists tarnishing.[9] When it first forms, this protective layer is only 1–2 nm thick but continues to slowly grow; reaching a thickness of 25 nm in four years.[16]
Titanium burns in air when heated to 1200 °C (2190 °F) and in pure oxygen when The Pourbaix diagram for titanium in pure water, heated to 610 °C (1130 °F) or higher, [15] perchloric acid or sodium hydroxide forming titanium dioxide.[7] As a result, the metal cannot be melted in open air as it burns before the melting point is reached, so melting is only possible in an inert atmosphere or in vacuum. It is also one of the few elements that burns in pure nitrogen gas (it burns at 800 °C or 1,472 °F and forms titanium nitride, which causes embrittlement).[17] Titanium is resistant to dilute sulfuric acid and hydrochloric acid, along with chlorine gas, chloride solutions, and most organic acids.[3] It is paramagnetic (weakly attracted to magnets) and has fairly low electrical and thermal conductivity.[9] Experiments have shown that natural titanium becomes radioactive after it is bombarded with deuterons, emitting mainly positrons and hard gamma rays.[3] When it is red hot the metal combines with oxygen, and when it reaches 550 °C (1022 °F) it combines with chlorine.[3] It also reacts with the other halogens and absorbs hydrogen.[4]
Compounds The +4 oxidation state dominates in titanium chemistry,[18] but compounds in the +3 oxidation state are also common.[19] Because of this high oxidation state, many titanium compounds have a high degree of covalent bonding. Star sapphires and rubies get their asterism from the titanium dioxide impurities present in them.[16] Titanates are compounds made with titanium dioxide. Barium titanate has piezoelectric properties, thus making it possible to use it as a transducer in the interconversion of sound and electricity.[7] Esters of titanium are formed by the reaction of alcohols and titanium tetrachloride and are used to waterproof fabrics.[7] Titanium nitride (TiN) is often used to coat cutting tools, such as drill bits.[20] It also finds use as a gold-colored decorative finish, and as a barrier metal in semiconductor fabrication.[21] Titanium tetrachloride (titanium(IV) chloride, TiCl4, sometimes called "Tickle") is a colorless liquid which is used as an intermediate in the manufacture of titanium dioxide for paint.[22] It is widely used in
TiN coated drill bit
Titanium
5
organic chemistry as a Lewis acid, for example in the Mukaiyama aldol condensation.[23] Titanium also forms a lower chloride, titanium(III) chloride (TiCl3), which is used as a reducing agent.[24] Titanocene dichloride is an important catalyst for carbon-carbon bond formation. Titanium isopropoxide is used for Sharpless epoxidation. Other compounds include titanium bromide (used in metallurgy, superalloys, and high-temperature electrical wiring and coatings) and titanium carbide (found in high-temperature cutting tools and coatings).[4]
Occurrence Producer
Thousands of tons
% of total
Australia
1291.0
30.6
South Africa
850.0
20.1
Canada
767.0
18.2
Norway
382.9
9.1
Ukraine
357.0
8.5
Other countries
573.1
13.6
Total world
4221.0
100.0 [25]
Source: 2003 production of titanium dioxide.
Due to rounding, values do not sum to 100%.
Titanium is always bonded to other elements in nature. It is the ninth-most abundant element in the Earth's crust (0.63% by mass)[26] and the seventh-most abundant metal. It is present in most igneous rocks and in sediments derived from them (as well as in living things and natural bodies of water).[3] [9] In fact, of the 801 types of igneous rocks analyzed by the United States Geological Survey, 784 contained titanium.[26] Its proportion in soils is approximately 0.5 to 1.5%.[26] It is widely distributed and occurs primarily in the minerals anatase, brookite, ilmenite, perovskite, rutile, titanite (sphene), as well in many iron ores.[16] Of these minerals, only rutile and ilmenite have any economic importance, yet even they are difficult to find in high concentrations. Significant titanium-bearing ilmenite deposits exist in western Australia, Canada, China, India, New Zealand, Norway, and Ukraine.[16] Large quantities of rutile are also mined in North America and South Africa and help contribute to the annual production of 90,000 tonnes of the metal and 4.3 million tonnes of titanium dioxide.[16] Total known reserves of titanium are estimated to exceed 600 million tonnes.[16] Titanium is contained in meteorites and has been detected in the sun and in M-type stars;[3] the coolest type of star with a surface temperature of 3200 °C (5790 °F).[27] Rocks brought back from the moon during the Apollo 17 mission are composed of 12.1% TiO2.[3] It is also found in coal ash, plants, and even the human body.
Titanium
Isotopes Naturally occurring titanium is composed of 5 stable isotopes: 46Ti, 47Ti, 48Ti, 49Ti, and 50 Ti, with 48Ti being the most abundant (73.8% natural abundance). Eleven radioisotopes have been characterized, with the most stable being 44Ti with a half-life of 63 years, 45Ti with a half-life of 184.8 minutes, 51Ti with a half-life of 5.76 minutes, and 52Ti with a half-life of 1.7 minutes. All of the remaining radioactive isotopes have half-lives that are less than 33 seconds and the majority of these have half-lives that are less than half a second.[8] The isotopes of titanium range in atomic weight from 39.99 u (40Ti) to 57.966 u (58Ti). The primary decay mode before the most abundant stable isotope, 48Ti, is electron capture and the primary mode after is beta emission. The primary decay products before 48Ti are element 21 (scandium) isotopes and the primary products after are element 23 (vanadium) isotopes.[8]
History Titanium was discovered included in a mineral in Cornwall, England, in 1791 by amateur geologist and pastor William Gregor, then vicar of Creed parish.[28] He recognized the presence of a new element in ilmenite[4] when he found black sand by a stream in the nearby parish of Manaccan and noticed the sand was attracted by a magnet.[28] Analysis of the sand determined the presence of two metal oxides; iron oxide (explaining the attraction to the magnet) and 45.25% of a white metallic oxide he could not identify.[26] Gregor, realizing that the unidentified oxide contained a metal that did not match the properties of any known element, reported his findings to the Royal Geological Society of Cornwall and in the German science journal Crell's Annalen.[28] Around the same time, Franz-Joseph Müller von Reichenstein produced a similar substance, but could not identify it.[4] The oxide was independently rediscovered in 1795 by German chemist Martin Heinrich Klaproth in rutile from Hungary.[28] Klaproth found that it contained a new element and named it for the Titans of Greek mythology.[27] After hearing about Gregor's earlier discovery, he obtained a sample of manaccanite and confirmed it contained titanium. The processes required to extract titanium from its various ores are laborious and costly; it is not possible to reduce in the normal manner, by heating in the presence of carbon, because Martin Heinrich Klaproth that produces titanium carbide.[28] Pure metallic titanium named titanium for the (99.9%) was first prepared in 1910 by Matthew A. Hunter at Titans of Greek mythology. Rensselaer Polytechnic Institute by heating TiCl4 with sodium in a steel bomb at 700–800 °C in the Hunter process.[3] Titanium metal was not used outside the laboratory until 1932 when William Justin Kroll proved that it could be produced by reducing titanium tetrachloride (TiCl4) with calcium.[29] Eight years later he refined this process by using magnesium and even sodium in what became known as the Kroll process.[29] Although research continues into more efficient and cheaper processes (e.g., FFC Cambridge), the Kroll process is still used for commercial production.[3] [4]
6
Titanium
7 Titanium of very high purity was made in small quantities when Anton Eduard van Arkel and Jan Hendrik de Boer discovered the iodide, or crystal bar, process in 1925, by reacting with iodine and decomposing the formed vapors over a hot filament to pure metal.[30]
In the 1950s and 1960s the Soviet Union pioneered the use of titanium in military and submarine applications (Alfa Class and Mike Class)[31] as part of programs related to the Cold War.[32] Starting in the early 1950s, titanium began to be used extensively for military aviation purposes, particularly in high-performance jets, starting with aircraft such as the F100 Super Sabre and Lockheed A-12. A titanium crystal bar made by the iodide process
In the USA, the Department of Defense realized the strategic importance of the metal[33] and supported early efforts of commercialization.[34] Throughout the period of the Cold War, titanium was considered a Strategic Material by the U.S. government, and a large stockpile of titanium sponge was maintained by the Defense National Stockpile Center, which was finally depleted in 2005.[35] Today, the world's largest producer, Russian-based VSMPO-Avisma, is estimated to account for about 29% of the world market share.[36] In 2006, the U.S. Defense Agency awarded $5.7 million to a two-company consortium to develop a new process for making titanium metal powder. Under heat and pressure, the powder can be used to create strong, lightweight items ranging from armor plating to components for the aerospace, transportation, and chemical processing industries.[37]
Production and fabrication The processing of titanium metal occurs in 4 major steps:[38] reduction of titanium ore into "sponge", a porous form; melting of sponge, or sponge plus a master alloy to form an ingot; primary fabrication, where an ingot is converted into general mill products such as billet, bar, plate, sheet, strip, and tube; and secondary fabrication of finished shapes from mill products. Because the metal reacts with oxygen at high Titanium (Mineral Concentrate) temperatures it cannot be produced by reduction of its dioxide.[10] Titanium metal is therefore produced commercially by the Kroll process, a complex and expensive batch process. (The relatively high market value of titanium is mainly due to its processing, which sacrifices another expensive metal, magnesium.[39] ) In the Kroll process, the oxide is first converted to chloride through carbochlorination, whereby chlorine gas is passed over red-hot rutile or ilmenite in the presence of carbon to make TiCl4. This is condensed and purified by fractional distillation and then reduced with 800 °C molten magnesium in an argon atmosphere.[7] A more recently developed method, the FFC Cambridge process,[40] may eventually replace the Kroll process. This method uses titanium dioxide powder (which is a refined form of rutile) as feedstock to make the end product which is either a powder or sponge. If mixed
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8
oxide powders are used, the product is an alloy manufactured at a much lower cost than the conventional multi-step melting process. The FFC Cambridge process may render titanium a less rare and expensive material for the aerospace industry and the luxury goods market, and could be seen in many products currently manufactured using aluminium and specialist grades of steel. Common titanium alloys are made by reduction. For example, cuprotitanium (rutile with copper added is reduced), ferrocarbon titanium (ilmenite reduced with coke in an electric furnace), and manganotitanium (rutile with manganese or manganese oxides) are reduced.[17] 2 FeTiO3 + 7 Cl2 + 6 C (900 °C) → 2 TiCl4 + 2 FeCl3 + 6 CO TiCl4 + 2 Mg (1100 °C) → 2 MgCl2 + Ti About 50 grades of titanium and titanium alloys are designated and currently used, although only a couple of dozen are readily available commercially.[41] The ASTM International recognizes 31 Grades of titanium metal and alloys, of which Grades 1 through 4 are commercially pure (unalloyed). These four are distinguished by their varying degrees of tensile strength, as a function of oxygen content, with Grade 1 being the most ductile (lowest tensile strength with an oxygen content of 0.18%), and Grade 4 the least (highest tensile strength with an oxygen content of 0.40%).[16] The remaining grades are alloys, each designed for specific purposes, be it ductility, strength, hardness, electrical resistivity, creep resistance, resistance to corrosion from specific media, or a combination thereof.[42] The grades covered by ASTM and other alloys are also produced to meet Aerospace and Military specifications (SAE-AMS, MIL-T), ISO standards, and country-specific specifications, as well as proprietary end-user specifications for aerospace, military, medical, and industrial applications.[43] In terms of fabrication, all welding of titanium must be done in an inert atmosphere of argon or helium in order to shield it from contamination with atmospheric gases such as oxygen, nitrogen, or hydrogen.[12] Contamination will cause a variety of conditions, such as embrittlement, which will reduce the integrity of the assembly welds and lead to joint failure. Commercially pure flat product (sheet, plate) can be formed readily, but processing must take into account the fact that the metal has a "memory" and tends to spring back. This is especially true of certain high-strength alloys.[44] [45] The metal can be machined using the same equipment and via the same processes as stainless steel.[12]
Applications Titanium is used in steel as an alloying element (ferro-titanium) to reduce grain size and as a deoxidizer, and in stainless steel to reduce carbon content.[9] Titanium is often alloyed with aluminium (to refine grain size), vanadium, copper (to harden), iron, manganese, molybdenum, and with other metals.[46] Applications for titanium mill products (sheet, plate, bar, wire, forgings, castings) can be found in industrial, aerospace, recreational, and emerging markets. Powdered titanium is used in pyrotechnics as a source of bright-burning particles.
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Pigments, additives and coatings About 95% of titanium ore extracted from the Earth is destined for refinement into titanium dioxide (TiO2), an intensely white permanent pigment used in paints, paper, toothpaste, and plastics.[47] It is also used in cement, in gemstones, as an optical opacifier in paper,[48] and a strengthening agent in graphite composite fishing rods and golf clubs. TiO2 powder is chemically inert, resists fading in sunlight, and is very opaque: this allows it to impart a Titanium dioxide is the most commonly used compound of titanium pure and brilliant white color to the brown or gray chemicals that form the majority of household plastics.[4] In nature, this compound is found in the minerals anatase, brookite, and rutile.[9] Paint made with titanium dioxide does well in severe temperatures, is somewhat self-cleaning, and stands up to marine environments.[4] Pure titanium dioxide has a very high index of refraction and an optical dispersion higher than diamond.[3] In addition to being a very important pigment, titanium dioxide is also used in sunscreens due to its ability to protect skin by itself.[10] Recently, it has been put to use in air purifiers (as a filter coating), or in film used to coat windows on buildings which when exposed to UV light (either solar or man-made) and moisture in the air produces reactive redox species like hydroxyl radicals that can purify the air or keep window surfaces clean.[49]
Aerospace and marine Due to their high tensile strength to density ratio,[7] high corrosion resistance,[3] and ability to withstand moderately high temperatures without creeping, titanium alloys are used in aircraft, armor plating, naval ships, spacecraft, and missiles.[3] [4] For these applications titanium alloyed with aluminium, vanadium, and other elements is used for a variety of components including critical structural parts, fire walls, landing gear, exhaust ducts (helicopters), and hydraulic systems. In fact, about two thirds of all titanium metal produced is used in aircraft engines and frames.[50] The SR-71 "Blackbird" was one of the first aircraft to make extensive use of titanium within its structure, paving the way for its use in modern military and commercial aircraft. An estimated 59 metric tons (130,000 pounds) are used in the Boeing 777, 45 in the Boeing 747, 18 in the Boeing 737, 32 in the Airbus A340, 18 in the Airbus A330, and 12 in the Airbus A320. The Airbus A380 may use 146 metric tons, including about 26 tons in the engines.[51] In engine applications, titanium is used for rotors, compressor blades, hydraulic system components, and nacelles. The titanium 6AL-4V alloy accounts for almost 50% of all alloys used in aircraft applications.[52] Due to its high corrosion resistance to sea water, titanium is used to make propeller shafts and rigging and in the heat exchangers of desalination plants;[3] in heater-chillers for salt water aquariums, fishing line and leader, and for divers' knives. Titanium is used to manufacture the housings and other components of ocean-deployed surveillance and monitoring devices for scientific and military use. The former Soviet Union developed techniques for making submarines largely out of titanium, which became both the fastest and deepest diving submarines of their time.[53]
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Industrial Welded titanium pipe and process equipment (heat exchangers, tanks, process vessels, valves) are used in the chemical and petrochemical industries primarily for corrosion resistance. Specific alloys are used in downhole and nickel hydrometallurgy applications due to their high strength titanium Beta C, corrosion resistance, or combination of both. The pulp and paper industry uses titanium in process equipment exposed to corrosive media such as sodium hypochlorite or wet chlorine gas (in the bleachery).[54] Other applications include: ultrasonic welding, wave soldering,[55] and sputtering targets.[56] Titanium tetrachloride (TiCl4), a colorless liquid, is important as an intermediate in the process of making TiO2 and is also used to produce the Ziegler-Natta catalyst, and is used to iridize glass and because it fumes strongly in moist air it is also used to make smoke screens.[10]
Consumer and architectural Titanium metal is used in automotive applications, particularly in automobile or motorcycle racing, where weight reduction is critical while maintaining high strength and rigidity.[57] The metal is generally too expensive to make it marketable to the general consumer market, other than high-end products, particularly for the racing/performance market. Late model Corvettes have been available with titanium exhausts.[58] Titanium is used in many sporting goods: tennis rackets, golf clubs, lacrosse stick shafts; cricket, hockey, lacrosse, and football helmet grills; and bicycle frames and components.[10] [59] Titanium alloys are also used in spectacle frames.[60] This results in a rather expensive, but highly durable and long lasting frame which is light in weight and causes no skin allergies. Many backpackers use titanium equipment, including cookware, eating utensils, lanterns, and tent stakes.[60] Though slightly more expensive than traditional steel or aluminium alternatives, these titanium products can be significantly lighter without compromising strength. Titanium is also favored for use by farriers, since it is lighter and more durable than steel when formed into horseshoes.[60] The Guggenheim Museum Bilbao is sheathed in titanium panels.
Because of its durability, titanium has become more popular for designer jewelry.[60] Its inertness makes it a good choice for those with allergies or those who will be wearing the jewelry in environments such as swimming pools. Titanium's durability, light weight, dent- and corrosion- resistance makes it useful in the production of watch cases.[60] A number of artists work with titanium to produce artworks such as sculptures, decorative objects and furniture. Titanium has occasionally been used in architectural applications: the 40 m (120 foot) memorial to Yuri Gagarin, the first man to travel in space, in Moscow, is made of titanium for the metal's attractive color and association with rocketry.[61] The Guggenheim Museum Bilbao and the Cerritos Millennium Library were the first buildings in Europe and North America, respectively, to be sheathed in titanium panels.[50] Other construction uses of titanium sheathing include the Frederic C. Hamilton Building in Denver, Colorado[62] and
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the 107 m (350 foot) Monument to the Conquerors of Space in Moscow.[63] Due to its superior strength and light weight when compared to other metals traditionally used in firearms (steel, stainless steel, and aluminium), and advances in metal-working techniques, the use of titanium has become more widespread in the manufacture of firearms. Primary uses include pistol frames and revolver cylinders. For these same reasons, it is also used in the body of laptop computers (for example, in Apple's PowerBook line).
Medical Because it is biocompatible (non-toxic and is not rejected by the body), titanium is used in a gamut of medical applications including surgical implements and implants, such as hip balls and sockets (joint replacement) that can stay in place for up to 20 years.[28] Titanium has the inherent property to osseointegrate, enabling use in dental implants that can remain in place for over 30 years.[28] This property is also useful for orthopedic implant applications.[28] This left lateral cephalametric
Since titanium is non-ferromagnetic, patients with radiograph shows a profile of the human skull. A fracture of the eye titanium implants can be safely examined with socket was repaired by stabilizing the magnetic resonance imaging (convenient for long-term fractured bones with small titanium implants). Preparing titanium for implantation in the plates and screws. body involves subjecting it to a high-temperature plasma arc which removes the surface atoms, exposing fresh titanium that is instantly oxidized.[28] Titanium is also used for the surgical instruments used in image-guided surgery, as well as wheelchairs, crutches, and any other products where high strength and low weight are desirable. Its inertness and ability to be attractively colored makes it a popular metal for use in body piercing.[64] Titanium may be anodized to produce various colors.[65]
Precautions Titanium is non-toxic even in large doses and does not play any natural role inside the human body.[27] An estimated 0.8 milligrams of titanium is ingested by humans each day but most passes through without being absorbed.[27] It does, however, have a tendency to bio-accumulate in tissues that contain silica. An unknown mechanism in plants may use titanium to stimulate the production of carbohydrates and encourage growth. This may explain why most plants contain about 1 part per million (ppm) of titanium, food plants have about 2 ppm, and horsetail and nettle contain up to 80 ppm.[27]
Nettle contains up to 80 parts per million of titanium
Titanium As a powder or in the form of metal shavings, titanium metal poses a significant fire hazard and, when heated in air, an explosion hazard.[66] Water and carbon dioxide-based methods to extinguish fires are ineffective on burning titanium; Class D dry powder fire fighting agents must be used instead.[4] When used in the production or handling of chlorine, care must be taken to use titanium only in locations where it will not be exposed to dry chlorine gas which can result in a titanium/chlorine fire.[67] A fire hazard exists even when titanium is used in wet chlorine due to possible unexpected drying brought about by extreme weather conditions. Titanium can catch fire when a fresh, non-oxidized surface comes in contact with liquid oxygen.[68] Such surfaces can appear when the oxidized surface is struck with a hard object, or when a mechanical strain causes the emergence of a crack. This poses the possible limitation for its use in liquid oxygen systems, such as those found in the aerospace industry.
See also • Titanium alloy • • • • • •
Titanium coating Titanium compounds Titanium in Africa Titanium minerals VSMPO-AVISMA Titanium Metals Corporation
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Titanium [15] Ignasi Puigdomenech, Hydra/Medusa Chemical Equilibrium Database and Plotting Software (2004) KTH Royal Institute of Technology, freely downloadable software at (http:/ / www. kemi. kth. se/ medusa/ ) [16] Emsley 2001, p. 453 [17] " Titanium (http:/ / encarta. msn. com/ encyclopedia_761569280/ Titanium. html)". Microsoft Encarta. 2005. . Retrieved 2006-12-29. [18] Greenwood 1997, p. 958 [19] Greenwood 1997, p. 970 [20] Truini, Joseph. " Drill Bits (http:/ / books. google. com/ books?id=Z-QDAAAAMBAJ& printsec=frontcover& source=gbs_summary_r& cad=0_0)". Popular Mechanics (Hearst Magazines) 165 (5): 91. ISSN 0032-4558 (http:/ / worldcat. org/ issn/ 0032-4558). . [21] Baliga, B. Jayant (2005). Silicon carbide power devices (http:/ / books. google. com/ books?id=LNLVwAzhN7EC& printsec=frontcover& source=gbs_summary_r& cad=0). World Scientific. pp. 91. ISBN 9812566058. . [22] Johnson, Richard W. (1998). The Handbook of Fluid Dynamics (http:/ / books. google. com/ books?id=JBTlucgGdegC). Springer. pp. 38–21. ISBN 3540646124. . [23] Coates, Robert M.; Paquette, Leo A. (2000). Handbook of Reagents for Organic Synthesis (http:/ / books. google. com/ books?id=xxYjJgupBSMC). John Wiley and Sons. pp. 93. ISBN 0470856254. . [24] Grimmett, M. Ross (1997). Imidazole and benzimidazole synthesis (http:/ / books. google. com/ books?id=jREKWf_hubkC). Academic Press. pp. 155. ISBN 0123031907. . [25] Cordellier, Serge; Didiot, Béatrice (2004). L'état du monde 2005: annuaire économique géopolitique mondial. Paris: La Découverte. [26] Barksdale 1968, p. 732 [27] Emsley 2001, p. 451 [28] Emsley 2001, p. 452 [29] Greenwood 1997, p. 955 [30] van Arkel, A. E.; de Boer, J. H. (1925). "Preparation of pure titanium, zirconium, hafnium, and thorium metal". Zeitschrift für anorganische und allgemeine Chemie 148: 345 – 50. [31] Yanko, Eugene; Omsk VTTV Arms Exhibition and Military Parade JSC (2006). " Submarines: general information (http:/ / warfare. ru/ ?lang=& linkid=1756& catid=243)". . Retrieved 2006-12-26. [32] Stainless Steel World (July/August 2001). " VSMPO Stronger Than Ever (http:/ / www. stainless-steel-world. net/ pdf/ ssw0107. pdf?issueID=30)". KCI Publishing B.V.. pp. 16–19. . Retrieved 2007-01-02. [33] NATIONAL MATERIALS ADVISORY BOARD, Commission on Engineering and Technical Systems (CETS), National Research Council (1983). Titanium: Past, Present, and Future (http:/ / books. nap. edu/ openbook. php?record_id=1712& page=R1). Washington, DC: national Academy Press. pp. R9. NMAB-392. . [34] " Titanium Metals Corporation. Answers.com. Encyclopedia of Company Histories, (http:/ / www. answers. com/ topic/ titanium-metals-corporation)". Answers Corporation. 2006. . Retrieved 2007-01-02. [35] Defense National Stockpile Center (2006) (PDF). Strategic and Critical Materials Report to the Congress. Operations under the Strategic and Critical Materials Stock Piling Act during the Period October 2004 through September 2005 (https:/ / www. dnsc. dla. mil/ . . \Uploads/ Materials/ admin_4-26-2006_14-19-33_SRC 2005 Ops Report Complete. pdf). United States Department of Defense. pp. § 3304. . [36] Bush, Jason (2006-02-15). " Boeing's Plan to Land Aeroflot (http:/ / www. businessweek. com/ technology/ content/ feb2006/ tc20060215_694672. htm?campaign_id=search)". BusinessWeek. . Retrieved 2006-12-29. [37] DuPont (2006-12-09). " U.S. Defense Agency Awards $5.7 Million to DuPont and MER Corporation for New Titanium Metal Powder Process (http:/ / www. highbeam. com/ doc/ 1G1-151246469. html)". . Retrieved 2009-08-01. [38] Matthew J. Donachie, Jr. (1988). TITANIUM: A Technical Guide. Metals Park, OH: ASM International. Chapter 4. ISBN 0871703092. [39] Barksdale 1968, p. 733 [40] Chen, George Zheng; Fray, Derek J.; Farthing, Tom W. (2000). " Direct electrochemical reduction of titanium dioxide to titanium in molten calcium chloride (http:/ / www. nature. com/ nature/ journal/ v407/ n6802/ full/ 407361a0. html)". Nature 407: 361–364. doi: 10.1038/35030069 (http:/ / dx. doi. org/ 10. 1038/ 35030069). . [41] Matthew J. Donachie, Jr. (1988). TITANIUM: A Technical Guide. Metals Park, OH: ASM International. pp. 16, Appendix J. ISBN 0871703092. [42] ASTM International (2006). Annual Book of ASTM Standards (Volume 02.04: Non-ferrous Metals). West Conshohocken, PA: ASTM International. section 2. ISBN 0-8031-4086-X. ASTM International (1998). Annual Book of ASTM Standards (Volume 13.01: Medical Devices; Emergency Medical Services). West Conshohocken, PA: ASTM International. sections 2 & 13. ISBN 0-8031-2452-X. [43] Matthew J. Donachie, Jr. (1988). TITANIUM: A Technical Guide. Metals Park, OH: ASM International. pgs.13–16, Appendices H and J. ISBN 0871703092.
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Titanium [44] AWS G2.4/G2.4M:2007 Guide for the Fusion Welding of Titanium and Titanium Alloys (http:/ / www. awspubs. com/ product_info. php?products_id=408). Miami: American Welding Society. 2006. . [45] Titanium Metals Corporation (1997). Titanium design and fabrication handbook for industrial applications (http:/ / www. timet. com/ design& fabframe. html). Dallas: Titanium Metals Corporation. . [46] Hampel, Clifford A. (1968). The Encyclopedia of the Chemical Elements. Van Nostrand Reinhold. pp. 738. ISBN 0442155980. [47] United States Geological Survey (2006-12-21). " USGS Minerals Information: Titanium (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ titanium/ )". . Retrieved 2006-12-29. [48] Smook, Gary A. (2002). Handbook for Pulp & Paper Technologists (3rd edition). Angus Wilde Publications. pp. 223. ISBN 0-9694628-5-9. [49] Stevens, Lisa; Lanning, John A.; Anderson, Larry G.; Jacoby, William A.; Chornet, Nicholas (June 14 – 18, 1998). " Photocatalytic Oxidation of Organic Pollutants Associated with Indoor Air Quality (http:/ / carbon. cudenver. edu/ ~landerso/ 98mp9b06. htm)". Air & Waste Management Association 91st Annual Meeting & Exhibition, San Diego. [50] Emsley 2001, p. 454 [51] Sevan, Vardan (2006-09-23). " Rosoboronexport controls titanium in Russia (http:/ / www. sevanco. net/ news/ full_story. php?id=1122)". Sevanco Strategic Consulting. . Retrieved 2006-12-26. [52] Matthew J. Donachie, Jr. (1988). TITANIUM: A Technical Guide. Metals Park, OH: ASM International. pp. 13,. ISBN 0871703092. [53] " GlobalSecurity (http:/ / www. globalsecurity. org/ military/ world/ russia/ 705. htm)". GlobalSecurity.org. April 2006. . Retrieved 2008-04-23. [54] Matthew J. Donachie, Jr. (1988). TITANIUM: A Technical Guide. Metals Park, OH: ASM International. 11–16. ISBN 0871703092. [55] E.W. Kleefisch, Editor (1981). Industrial Application of Titanium and Zirconium (http:/ / www. astm. org/ cgi-bin/ SoftCart. exe/ BOOKSTORE/ PUBS/ 943. htm?E+ mystore). West Conshohocken, PA: ASTM International. ISBN 0803107455. . [56] Rointan F. Bunshah, Editor (2001). Handbook of Hard Coatings (http:/ / books. google. com/ books?id=daamnz8el2sC& pg=PA413). Norwich, NY: William Andrew Inc.. pp. Ch. 8. ISBN 0815514387. . [57] Bell, Tom; et al. (2001). Heat Treating (http:/ / books. google. com/ books?id=4F1zYT4FHyMC). Proceedings of the 20th Conference, 9-12 October 2000. ASM International. pp. 141. ISBN 0871707276. . [58] National Corvette Museum (2006). " Titanium Exhausts (http:/ / www. iglou. com/ corvette/ specs/ 2001/ exhaust. htm)". . Retrieved 2006-12-26. [59] Davis, Joseph R. (1998). Metals Handbook (http:/ / books. google. com/ books?id=IpEnvBtSfPQC). ASM International. pp. 584. ISBN 0871706547. . [60] Donachie, Matthew J. (2000). Titanium: A Technical Guide. ASM International. pp. 11, 255. ISBN 0871706865. [61] " Yuri Gagarin (http:/ / encarta. msn. com/ encyclopedia_761571506/ Gagarin_Yuri_Alekseyevich. html)". Microsoft Encarta. 2006. . Retrieved 2006-12-26. [62] " Denver Art Museum, Frederic C. Hamilton Building (http:/ / www. designbuild-network. com/ projects/ dam/ )". SPG Media. 2006. . Retrieved 2006-12-26. [63] Gruntman, Mike (AIAA). Blazing the Trail: The Early History of Spacecraft and Rocketry (http:/ / books. google. com/ books?id=2XY9KXxF8OEC). Reston, VA: American Institute of Aeronautics and Astronautics. pp. 457. ISBN 156347705X. . [64] " Body Piercing Safety (http:/ / www. doctorgoodskin. com/ tp/ bodypiercing/ )". . Retrieved 2009-08-01. [65] Alwitt, Robert S. (2002). " Electrochemistry Encyclopedia (http:/ / electrochem. cwru. edu/ ed/ encycl/ art-a02-anodizing. htm)". . Retrieved 2006-12-30. [66] Cotell, Catherine Mary; Sprague, J. A.; Smidt, F. A. (1994). ASM Handbook: Surface Engineering (http:/ / books. google. com/ books?id=RGtsPjqUwy0C) (10th ed.). ASM International. pp. 836. ISBN 087170384X. . [67] Compressed Gas Association (1999). Handbook of compressed gases (http:/ / books. google. com/ books?id=WSLULtCG9JgC) (4th ed.). Springer. pp. 323. ISBN 0412782308. . [68] Solomon, Robert E. (2002). Fire and Life Safety Inspection Manual (http:/ / books. google. com/ books?id=2fHsoobsCNwC). National Fire Prevention Association (8th ed.). Jones & Bartlett Publishers. pp. 45. ISBN 0877654727. .
• Flower, Harvey M. (2000). " Materials Science: A moving oxygen story (http:/ / www. nature. com/ nature/ journal/ v407/ n6802/ full/ 407305a0. html)". Nature 407: 305–306. doi: 10.1038/35030266 (http:/ / dx. doi. org/ 10. 1038/ 35030266). http:/ / www. nature. com/ nature/ journal/ v407/ n6802/ full/ 407305a0. html.
14
Titanium • Stwertka, Albert (1998). Guide to the Elements (Revised Edition). Oxford: Oxford University Press. ISBN 0-19-508083-1. • Winter, Mark (2006). " Chemistry: Periodic table: Titanium (http:/ / www. webelements. com/ webelements/ elements/ text/ Ti/ index. html)". WebElements. http:/ / www. webelements. com/ webelements/ elements/ text/ Ti/ index. html. Retrieved 2006-12-10.
Bibliography • Barksdale, Jelks (1968). "Titanium". in Clifford A. Hampel (editor). The Encyclopedia of the Chemical Elements. New York: Reinhold Book Corporation. pp. 732–738. LCCN 68-29938. • CRC contributors (2006). David R. Lide (editor). ed. Handbook of Chemistry and Physics (87th ed.). Boca Raton, Florida: CRC Press, Taylor & Francis Group. ISBN 0-8493-0487-3. • Emsley, John (2001). "Titanium". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. pp. 457–456. ISBN 0198503407. • Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. ISBN 0-7506-3365-4.
External links • A Cleaner, Cheaper Route to Titanium (http:/ / www. techreview. com/ read_article. aspx?id=16963& ch=nanotech) • International Titanium Association (http:/ / www. titanium. org) • Metallurgy of Titanium and its Alloys, Cambridge University (http:/ / www. msm. cam. ac. uk/ phase-trans/ 2003/ titanium. movies/ titanium. html) • World Production of Titanium Concentrates, by Country (http:/ / www. indexmundi. com/ en/ commodities/ minerals/ titanium/ titanium_table15. html) • Truth in Sparks: Titanium or Plain Ol' Steel? (http:/ / www. popsci. com/ popsci/ how20/ 85f145ef7d2f6110vgnvcm1000004eecbccdrcrd. html) Popular Science Magazine
15
Article Sources and Contributors
Article Sources and Contributors Titanium Source: http://en.wikipedia.org/w/index.php?oldid=307495308 Contributors: .:Ajvol:., 129.186.19.xxx, 334a, 56, A. Parrot, ABF, AUG, Abce2, Abigail-II, Abrech, Acalamari, Acroterion, Adashiel, Addshore, AdjustShift, Adrian, Ahoerstemeier, Ajnin, AlanCatchpole, Alansohn, Alchemist-hp, Aldaron, Aldis90, AlexanderM, Alexfusco5, AlexiusHoratius, Alexy527, Algebraist, Alison, Alphachimp, Anclation, AndonicO, AndySnow, Angela, Anonymous101, Antandrus, Anthere, Antonio Lopez, Antonrojo, Arado, Arakunem, Archimerged, Aremith, Arteitle, Ashley Payne, AssistantX, AubreyEllenShomo, Axlq, Azwatchdog, Backslash Forwardslash, Bandy, Bcdm, Bcorr, Beetstra, Ben 2012, Ben.c.roberts, Benbest, Benc, Bendy d, Bfesser, Bhadani, BillFlis, Bingkris, Bjmcculloch, Bkell, Blazenbbw, Blazotron, Bllrby6, BlueEarth, Bobblewik, Bobo192, Boccobrock, Bogey97, Bongwarrior, Bookofjude, BorgQueen, Boris Allen, Bostwickenator, Bovineone, Bowdyyz69, Brf, Brian Huffman, Brickc1, Brighterorange, Bryan Derksen, Bubbha, Bushellman, Bwrs, CASfan, CP\M, Cadmium, Call me Bubba, Caltas, CambridgeBayWeather, Can't sleep, clown will eat me, 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Image Sources, Licenses and Contributors file:hexagonal.svg Source: http://en.wikipedia.org/w/index.php?title=File:Hexagonal.svg License: BSD Contributors: Original uploader was Danieljamesscott at en.wikipedia file:Electron shell 022 Titanium.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_022_Titanium.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 file:Ti,22.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Ti,22.jpg License: GNU Free Documentation License Contributors: Eusebius, Paginazero, Saperaud, Silverhill File:Titanium in water porbiax diagram.png Source: http://en.wikipedia.org/w/index.php?title=File:Titanium_in_water_porbiax_diagram.png License: Public Domain Contributors: Original uploader was Cadmium at en.wikipedia File:Titanium nitride coating.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Titanium_nitride_coating.jpg License: unknown Contributors: Peter Binter Original uploader was Binter at de.wikipedia File:Martin Heinrich Klaproth.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Martin_Heinrich_Klaproth.jpg License: Public Domain Contributors: Gabor, Mu, Stern, Väsk, Андрей Романенко File:Titan-crystal bar.JPG Source: http://en.wikipedia.org/w/index.php?title=File:Titan-crystal_bar.JPG License: unknown Contributors: Alchemist-hp ( pse-mendelejew.de) File:TitaniumUSGOV.jpg Source: http://en.wikipedia.org/w/index.php?title=File:TitaniumUSGOV.jpg License: Public Domain Contributors: Ra'ike, Saperaud File:Titanium(IV) oxide.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Titanium(IV)_oxide.jpg License: unknown Contributors: Original uploader was Walkerma at en.wikipedia File:GuggenheimBilbao.jpg Source: http://en.wikipedia.org/w/index.php?title=File:GuggenheimBilbao.jpg License: GNU Free Documentation License Contributors: E. Goergen aka File:Lateralcephplated.JPG Source: http://en.wikipedia.org/w/index.php?title=File:Lateralcephplated.JPG License: Public Domain Contributors: DRosenbach File:Kopiva.JPG Source: http://en.wikipedia.org/w/index.php?title=File:Kopiva.JPG License: unknown Contributors: Pokrajac
16
License
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
17
Vanadium
1
Vanadium titanium ← vanadium → chromium ↑ V ↓ Nb
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
23V Periodic table
Appearance blue-silver-grey metal
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
2101
2289
2523
2814
3187
3679
Vanadium
2
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1414 kJ·mol−1 3rd: 2830 kJ·mol−1Atomic radiusCovalent radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusPoisson ratioMohs hardnessCAS registry number Most stable isotopes Main article: Isotopes of vanadium iso
N.A.
half-life
DM
DE (MeV)
DP
48
syn
15.9735 d
ε+β+
4.0123
48
49
syn
330 d
ε
0.6019
49
50
0.25%
1.5×1017y
ε
2.2083
50
β−
1.0369
50
V V V
51
V
99.75%
Ti Ti Ti Cr
51
V is stable with 28 neutron
vanadium, V, 23 transition metal5, 4, d50.9415(1) g·mol−1 [Ar] 3d3 4s2 2, 8, 11, 2 (Image) solid 6.0 g·cm−3 5.5 g·cm−3 2183 K,1910 °C,3470 °F 3680 K,3407 °C,6165 °F 21.5 kJ·mol−1 459 kJ·mol−1 (25 °C) 24.89 J·mol−1·K−15, 4, 3, 2, 1, -1 (amphoteric oxide) 1.63 (Pauling scale) 1st: 650.9 kJ·mol−1134 pm153±8 pm body-centered cubic paramagnetic (20 °C) 197 nΩ·m (300 K) 30.7 W·m−1·K−1 (25 °C) 8.4 µm·m−1·K−1 (20 °C) 4560 m/s 128 GPa 47 GPa 160 GPa 0.37 6.7 7440-62-2 Vanadium (pronounced /vəˈneɪdiəm/) is the chemical element with the symbol V and atomic number 23. It is a soft, silvery grey, ductile transition metal. The formation of an oxide layer stabilizes the metal against oxidation. Andrés Manuel del Río discovered vanadium in 1801 by analyzing the mineral vanadinite, and named it erythronium. Four years later, however, he was convinced by other scientists that erythronium was identical to chromium. The element was rediscovered in 1831 by Nils Gabriel Sefström, who named it vanadium after the Norse goddess of beauty and fertility, Vanadis (Freya). Both names were attributed to the wide range of colors found in vanadium compounds. The element occurs naturally in about 65 different minerals and in fossil fuel deposits. It is produced in China and Russia from steel smelter slag; other countries produce it either from the flue dust of heavy oil, or as a byproduct of uranium mining. It is mainly used to produce specialty steel alloys such as high speed tool steels. The compound vanadium pentoxide is used as a catalyst for the production of sulfuric acid. Vanadium is found in many organisms, and is used by some life forms as an active center of enzymes.
History Vanadium was originally discovered by Andrés Manuel del Río, a Spanish-born Mexican mineralogist, in 1801. Del Río extracted the element from a sample of Mexican "brown lead" ore, later named vanadinite. He found that its salts exhibit a wide variety of colors, and as a result he named the element panchromium (Greek: all colors). Later, Del Río renamed the element erythronium as most of its salts turned red upon heating. In 1805, the French chemist Hippolyte Victor Collet-Descotils, backed by del Río's friend, Baron Alexander von Humboldt, incorrectly declared that del Río's new element was only an impure sample of chromium. Del Río accepted the Collet-Descotils' statement, and retracted his claim.[1]
Vanadium In 1831, the Swedish chemist, Nils Gabriel Sefström, rediscovered the element in a new oxide he found while working with iron ores. Later that same year, Friedrich Wöhler confirmed del Río's earlier work.[2] Sefström choose a name beginning with V, which had not been assigned to any element yet. He called the element vanadium after Vanadis (another name for Freya, the Norse goddess of beauty and fertility), because of the many beautifully colored chemical compounds it produces.[2] In 1831, the geologist George William Featherstonhaugh suggested that vanadium should be renamed "rionium" after del Río, but this suggestion was not followed.[3] The isolation of vanadium metal proved difficult. In 1831, Berzelius reported the production of the metal, but Henry Enfield Roscoe showed that Berzelius had in fact produced the nitride, vanadium nitride (VN). Roscoe eventually produced the metal in 1867 by reduction of vanadium(III) chloride, VCl3, with hydrogen.[4] In 1927, pure vanadium was produced by reducing vanadium pentoxide with calcium.[5] The first large scale industrial use of vanadium in steels was found in the chassis of the Ford Model T, inspired by 1910 Model T French race cars. Vanadium steel allowed for reduced weight while simultaneously increasing tensile strength.[6]
Characteristics Vanadium is a soft, ductile, silver-grey metal. It has good resistance to corrosion and it is stable against alkalis, sulfuric and hydrochloric acids.[7] It is oxidized in air at about 933 K (660 °C, 1220 °F), although an oxide layer forms even at room temperature.
Isotopes Naturally occurring vanadium is composed of one stable isotope 51V and one radioactive isotope 50V. The latter has a half-life of 1.5×1017 years and a natural abundance 0.25%. 51V has a nuclear spin of 7/2 which is useful for NMR spectroscopy.[8] A number of 24 artificial radioisotopes have been characterized, ranging in mass number from 40 to 65. The most stable of these isotopes are 49V with a half-life of 330 days, and 48V with a half-life of 15.9735 days. All of the remaining radioactive isotopes have half-lives shorter than an hour, most of which are below 10 seconds. At least 4 isotopes have metastable excited states.[8] Electron capture is the main decay mode for isotopes lighter than the 51V. For the heavier ones, the most common mode is beta decay. The electron capture reactions lead to the formation of element 22 (titanium) isotopes, while for beta decay, it leads to element 24 (chromium) isotopes.
3
Vanadium
4
Chemistry and compounds The chemistry of vanadium is noteworthy for the accessibility of four adjacent oxidation states. The common oxidation states of vanadium are +2 (lilac), +3 (green), +4 (blue) and +5 (yellow). Vanadium(II) compounds are reducing agents, and vanadium(V) compounds are oxidizing agents. Vanadium(IV) compounds often exist as vanadyl derivatives which contain the VO2+ center.[7]
Vanadium(V) oxide is a catalyst in the Contact process for producing sulfuric acid
Ammonium vanadate(V) (NH VO ) can be successively 4
3
reduced with elemental zinc to obtain the different colors of vanadium in these four oxidation states. Lower oxidation states occur in compounds such as V(CO)6,[V(CO)6]- and substituted derivatives.[7] The vanadium redox battery utilizes these oxidation states; conversions of these oxidation states is illustrated by the reduction of a strongly acidic solution of a vanadium(V) compound with zinc dust. The initial yellow color characteristic of the vanadate ion, VO4, is replaced by the blue color of [VO(H2O)5]2+, followed by the green color of [V(H2O)6]3+ and then violet, due to [V(H2O)6]2+.[7] The most commercially important compound is vanadium pentoxide, which is used as a catalyst for the production of sulfuric acid.[7] This compound oxidizes sulfur dioxide (SO2) to the trioxide (SO3). In this redox reaction, sulfur is oxidized from +4 to +6, and vanadium is reduced from +5 to +3: V2O5 + 2 SO2 → V2O3 + 2 SO3 The catalyst is regenerated by oxidation with air: V2O3 + O2 → V2O5
Metavanadate chains
Vanadium
5
Oxy and oxo compounds
The Pourbaix diagram for vanadium in [9] water.
The oxyanion chemistry of vanadium(V) is complex. The vanadate ion, VO4, is present in dilute solutions at high pH. On acidification, HVO4 and H2VO4 are formed, analogous to HPO4 and H2PO4. The acid dissociation constants for the vanadium and phosphorus series are remarkably similar. In more concentrated solutions many polyvanadates are formed. Chains, rings and clusters involving tetrahedral vanadium, analogous to the polyphosphates, are known. In addition, clusters such as the decavanadates V10O28 and HV10O28, which predominate at pH 4-6, are formed in which compound [7] is octahedral about vanadium.
The correspondence between vanadate and phosphate chemistry can be attributed to the similarity in size and charge of phosphorus(V) and vanadium(V). Orthovanadate VO4 is used in protein crystallography[10] to study the biochemistry of phosphate.[11] Halide compounds Several halides are known for oxidation states +2, +3 and +4. VCl4 is the most important commercially. This liquid is mainly used as a catalyst for polymerization of dienes. Coordination compounds Vanadium's early position in the transition metal series lead to three rather unusual features of the coordination chemistry of vanadium. Firstly, metallic vanadium has the electronic configuration [Ar]4s23d3, so compounds of vanadium are relatively electron-poor. Consequently, most binary compounds are Lewis acids (electron pair acceptors); examples are all the halides forming octahedral adducts with the formula VXnL6-n (X A ball-and-stick model of VO(acac)2 = halide; L = other ligand). Secondly, the vanadium ion is rather large and can achieve coordination numbers higher than 6, as is the case in [V(CN)7]4−. Thirdly, the vanadyl ion, VO2+, is featured in many complexes of vanadium(IV) such as vanadyl acetylacetonate (V(=O)(acac)2). In this complex, the vanadium is 5-coordinate, square pyramidal, meaning that a sixth ligand, such as pyridine, may be attached, though the association constant of this process is small. Many 5-coordinate vanadyl complexes have a trigonal bypyramidal geometry, such as VOCl2(NMe3)2.[12] Organometallic compounds Organometallic chemistry of vanadium is well developed, but organometallic compounds are of minor commercial significance. Vanadocene dichloride is a versatile starting reagent and even finds minor applications in organic chemistry.[13] Vanadium carbonyl, V(CO)6, is a rare example of a metal carbonyl containing an unpaired electron, but which exists without dimerization. The addition of an electron yields V(CO)6 (isoelectronic with Cr(CO)6), which may be further reduced with sodium in liquid ammonia to yield V(CO)6 (isoelectronic with
Vanadium
6
Fe(CO)5).[14]
[15]
Occurrence
Vanadinite
Metallic vanadium is not found in nature, but is known to exist in about 65 different minerals. Economically significant examples include patronite (VS4),[16] vanadinite (Pb5(VO4)3Cl), and carnotite (K2(UO2)2(VO4)2·3H2O). Much of the world's vanadium production is sourced from vanadium-bearing magnetite found in ultramafic gabbro bodies. Vanadium is mined mostly in South Africa, north-western China, and eastern Russia. In 2007 these three countries mined more than 95 % of the 58,600 tonnes of
produced vanadium.[17] Vanadium is also present in bauxite and in fossil fuel deposits such as crude oil, coal, oil shale and tar sands. In crude oil, concentrations up to 1200 ppm have been reported. When such oil products are burned, the traces of vanadium may initiate corrosion in motors and boilers.[18] An estimated 110,000 tonnes of vanadium per year are released into the atmosphere by burning fossil fuels.[19] Vanadium has also been detected spectroscopically in light from the Sun and some other stars.[20]
Production Most vanadium is used as ferrovanadium as an additive to improve steels. Ferrovanadium is produced directly by reducing a mixture of vanadium oxide, iron oxides and iron in an electric furnace. Vanadium-bearing magnetite iron ore is the main source for the production of vanadium.[21] The vanadium ends up in pig iron produced from vanadium bearing magnetite. During steel production, oxygen is blown into the pig iron, oxidizing the carbon and most of the other impurities, forming slag. Depending on the used ore, the slag contains up to 25% of vanadium.[21] Vanadium metal is obtained via a multistep process that begins with the roasting of crushed ore with NaCl or Na2CO3 at about 850 °C to give sodium metavanadate (NaVO3). An aqueous extract of this solid is acidified to give "red cake", a polyvanadate salt, which is reduced with calcium metal. As an alternative for small scale production, vanadium pentoxide is reduced with hydrogen or magnesium. Many other methods are also in use, in all of which vanadium is produced as a byproduct of other processes.[21] Purification of vanadium is possible by the crystal bar process developed by Anton Eduard van Arkel and Jan Hendrik de Boer in 1925. It involves the formation of the metal iodide, in this example vanadium(III) iodide, and the subsequent decomposition to yield pure metal.[22] 2 V + 3 I2
2 VI3
Vanadium
7
Applications Alloys
Tool made from vanadium steel
Approximately 85% of vanadium produced is used as ferrovanadium or as a steel additive.[21] The considerable increase of strength in steel containing small amounts of vanadium was discovered in the beginning of 20th century,[23] and from that time vanadium steel was used for applications in axles, bicycle frames, crankshafts, gears, and other critical components. Vanadium forms stable nitrides and carbides, resulting in a significant increase in the strength of the steel. There are two groups of vanadium containing steel alloy groups. Vanadium high-carbon steel alloys containing 0.15 to 0.25 percent vanadium and high speed tool steels (HSS) with a vanadium content ranges from 1 % to 5 %. For high speed tool steels, a hardness above HRC 60 can be achieved. HSS steel is used in surgical instruments and tools.[24] Vanadium stabilizes the beta form of titanium and increases the strength and temperature stability of titanium. Mixed with aluminium in titanium alloys it is used in jet engines and high-speed airframes. One of the common alloys is Ti 6Al 4V a titanium alloy with 6% aluminium and 4% vanadium.[25]
Other uses Vanadium
is
compatible
with
iron
and
titanium,
therefore vanadium foil is used in cladding titanium to steel.[26] The moderate thermal neutron-capture cross-section and the short half-life of the isotopes produced by neutron capture makes vanadium a suitable material for the inner structure of a fusion reactor.[27] [28] Several vanadium alloys show superconducting behaviour. The first A15 phase superconductor was a vanadium compound, V3Si, which was discovered in 1952.[29] Vanadium-gallium tape is used in superconducting magnets (17.5 teslas or 175,000 gauss). The structure of the superconducting A15 phase of V3Ga is similar to that of the more common Nb3Sn and Nb3Ti.[30] Vanadium
The most common oxide of vanadium Vanadium pentoxide, V2O5, is used as a catalyst in manufacturing sulfuric acid by the contact process[31] and as an oxidizer in maleic anhydride production.[32] Vanadium pentoxide is also used in making ceramics.[33] Another oxide of vanadium, vanadium dioxide VO2, is used in the production of glass coatings, which blocks infrared radiation (and not visible light) at a specific temperature.[34] Vanadium oxide can be used to induce color centers in corundum to create simulated alexandrite jewelry, although alexandrite in nature is a chrysoberyl.[35] The possibility to use vanadium redox couples in both half-cells, thereby eliminating the problem of cross contamination by
Vanadium
8
diffusion of ions across the membrane is the advantage of vanadium redox rechargeable batteries.[36] Vanadate can be used for protecting steel against rust and corrosion by electrochemical conversion coating.[37] Lithium vanadium oxide has been proposed for use as a high energy density anode for lithium ion batteries, at 745 Wh/l when paired with a lithium cobalt oxide cathode.[38] It has been proposed by some researchers that a small amount, 40 to 270 ppm, of vanadium in Wootz steel and Damascus steel, significantly improves the strength of the material, although it is unclear what the source of the vanadium was.[39]
Biological role
Ascidiacea contain vanadium.
Amanita muscaria contains amavadin.
Vanadium plays a very limited role in biology. A vanadium-containing nitrogenase is used by some nitrogen-fixing micro-organisms. Vanadium is essential to ascidians or sea squirts in vanadium chromagen proteins. The concentration of vanadium in their blood is more than 100 times higher than the concentration of vanadium in the seawater around them. Rats and chickens are also known to require vanadium in very small amounts and deficiencies result in reduced growth and impaired reproduction.[40] Vanadium is a relatively controversial dietary supplement, primarily for increasing insulin sensitivity[41] and body-building. Whether it works for the latter purpose has not been proven, and there is some evidence that athletes who take it are merely experiencing a placebo effect.[42] Vanadyl sulfate may improve glucose control in people with type 2 diabetes.[43] [44] [45] [46] [47] . In addition, decavanadate and oxovanadates are species that potentially have many biological activities and that have been successfully used as tools in the comprehension of several biochemical processes.[48]
Ten percent of the blood cell pigment of the sea cucumber is vanadium. Just as the horseshoe crab has blue blood due to copper in hemocyanin, and land animals have red blood from the iron in hemoglobin, the blood of the sea cucumber is yellow because of the vanadium in the vanabin pigment.[49] Nonetheless, there is no evidence that vanabins carry oxygen, in contrast to hemoglobin and hemocyanin.[50] Several species of macrofungi, namely Amanita muscaria and related species, accumulate vanadium (up to 500 mg/kg in dry weight). Vanadium is present in the coordination complex, amavadin,[51] in fungal fruit-bodies. However, the biological importance of the accumulation process is unknown.[52] [53]
Vanadium
Safety All vanadium compounds should be considered to be toxic. Tetravalent VOSO4 has been reported to be over 5 times more toxic than trivalent V2O3.[54] The Occupational Safety and Health Administration (OSHA) has set an exposure limit of 0.05 mg/m3 for vanadium pentoxide dust and 0.1 mg/m3 for vanadium pentoxide fumes in workplace air for an 8-hour workday, 40-hour work week.[55] The National Institute for Occupational Safety and Health (NIOSH) has recommended that 35 mg/m3 of vanadium be considered immediately dangerous to life and health. This is the exposure level of a chemical that is likely to cause permanent health problems or death.[55] Vanadium compounds are poorly absorbed through the gastrointestinal system. Inhalation exposures to vanadium and vanadium compounds result primarily in adverse effects on the respiratory system.[56] [57] [58] Quantitative data are, however, insufficient to derive a subchronic or chronic inhalation reference dose. Other effects have been reported after oral or inhalation exposures on blood parameters,[59] [60] on liver,[61] on neurological development in rats,[62] and other organs.[63] There is little evidence that vanadium or vanadium compounds are reproductive toxins or teratogens. Vanadium pentoxide was reported to be carcinogenic in male rats and male and female mice by inhalation in an NTP study,[64] although the interpretation of the results has recently been disputed.[65] Vanadium has not been classified as to carcinogenicity by the U.S. EPA.[66]
External links • The periodic table of videos [67] videos of the chemistry of the elements • WebElements.com – Vanadium [68] • ATSDR – ToxFAQs: Vanadium [69]
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Corrosion 60 (3): 284–296. [38] Kariatsumari, Koji (February 2008). " Li-Ion Rechargeable Batteries Made Safer (http:/ / techon. nikkeibp. co. jp/ article/ HONSHI/ 20080129/ 146549/ )". Nikkei Business Publications, Inc.. . Retrieved 10-12-2008. [39] Verhoeven, J. D.; Pendray, A. H.; Dauksch, W. E. (1998). "The key role of impurities in ancient damascus steel blades". Journal of the Minerals, Metals and Materials Society 50 (9): 58–64. doi: 10.1007/s11837-998-0419-y (http:/ / dx. doi. org/ 10. 1007/ s11837-998-0419-y). [40] Schwarz, Klaus; Milne, David B. (1971). " Growth Effects of Vanadium in the Rat (http:/ / www. jstor. org/ stable/ 1731776)". Science 174 (4007): 426–428. doi: 10.1126/science.174.4007.426 (http:/ / dx. doi. org/ 10. 1126/ science. 174. 4007. 426). PMID 5112000. . [41] Yeh, Gloria Y.; Eisenberg, David M.; Kaptchuk, Ted J.; Phillips, Russell S. (2003). " Systematic Review of Herbs and Dietary Supplements for Glycemic Control in Diabetes (http:/ / care. diabetesjournals. org/ cgi/ content/ full/ 26/ 4/ 1277)". Diabetes Care 26: 1277–1294. doi: 10.2337/diacare.26.4.1277 (http:/ / dx. doi. org/ 10. 2337/ diacare. 26. 4. 1277). PMID 12663610. . [42] Talbott,, Shawn M.; Hughes, Kerry (2007). " Vanadium (http:/ / books. google. com/ books?id=hV2_TdmoDo8C& pg=PA419)". The Health Professional's Guide to Dietary Supplements. Lippincott Williams & Wilkins. pp. 419–422. ISBN 9780781746724. . [43] Halberstam, M, et al. (1996). "Oral vanadyl sulfate improves insulin sensitivity in NIDDM but not in obese nondiabetic subjects.". Diabetes 45: 659–66. doi: 10.2337/diabetes.45.5.659 (http:/ / dx. doi. org/ 10. 2337/ diabetes. 45. 5. 659). PMID 8621019. [44] Boden, G, et al. (1996;). "Effects of vanadyl sulfate on carbohydrate and lipid metabolism in patients with non-insulin dependent diabetes mellitus.". Metabolism 45: 1130–5. doi: 10.1016/S0026-0495(96)90013-X (http:/ / dx. doi. org/ 10. 1016/ S0026-0495(96)90013-X). [45] Goldfine, AB, et al. (2000). "Metabolic effects of vanadyl sulfate in humans with non-insulin-dependent diabetes mellitus: in vivo and in vitro studies.". Metabolism 49: 400–10. doi: 10.1016/S0026-0495(00)90418-9 (http:/ / dx. doi. org/ 10. 1016/ S0026-0495(00)90418-9). [46] Badmaev, V, et al. (1999). "Vanadium: a review of its potential role in the fight against diabetes.". Altern Complement Med. 5: 273–291. doi: 10.1089/acm.1999.5.273 (http:/ / dx. doi. org/ 10. 1089/ acm. 1999. 5. 273). [47] Goldwaser, I, et al. (1999). "L-Glutamic Acid gamma -Monohydroxamate. A Potentiator of Vanadium-Evoked Glucose Metabolism in vitro and in vivo". J Biol Chem 274: 26617–26624. doi: 10.1074/jbc.274.37.26617 (http:/ / dx. doi. org/ 10. 1074/ jbc. 274. 37. 26617). PMID 10473627. [48] Aureliano, Manuel; Crans, Debbie C. (2009). "Decavanadate and oxovanadates: Oxometalates with many biological activities". Journal Inorganic Biochemistry 103: 536–546. doi: 10.1016/j.jinorgbio.2008.11010 (http:/ / dx. doi. org/ 10. 1016/ j. jinorgbio. 2008. 11010). [49] Natkin, Michael (2007). " Blood Color (http:/ / www. soak. com/ topic/ sciencefacts/ article/ tshow/ 98556/ blood+ color)". Science Facts. Soak (Source Of All Knowledge). . Retrieved 2007-11-16. [50] Rehder, Dieter (1992). "Structure and function of vanadium compounds in living organisms". BioMetals 5 (1): 3–12. doi: 10.1007/BF01079691 (http:/ / dx. doi. org/ 10. 1007/ BF01079691). [51] Kneifel, Helmut; Bayer, Ernst (1997). "Determination of the Structure of the Vanadium Compound, Amavadine, from Fly Agaric". Angewandte Chemie International Edition in English 12 (6): 508. ISSN 10.1002/anie.197305081 (http:/ / worldcat. org/ issn/ 10. 1002/ anie. 197305081). [52] Falandysz, J.; Kunito, T., Kubota, R.; Lipka, K.; Mazur, A.; Falandysz, Justyna J.; Tanabe, S. (2007). "Selected elements in fly agaric Amanita muscaria". Journal of Environmental Science and Health, Part A 42 (11): 1615–1623. doi: 10.1080/10934520701517853 (http:/ / dx. doi. org/ 10. 1080/ 10934520701517853). [53] Berry, Robert E.; Armstrong, Elaine M.; Beddoes, Roy L.; Collison, David; Ertok, Nigar; Helliwell, Madeleine; Garner, David (1999). "The Structural Characterization of Amavadin". Angew. Chem. Int. Ed. 38 (6): 795–797. doi: 10.1002/(SICI)1521-3773(19990315)38:6<795::AID-ANIE795>3.0.CO;2-7 (http:/ / dx. doi. org/ 10. 1002/ (SICI)1521-3773(19990315)38:6<795::AID-ANIE795>3. 0. CO;2-7). [54] Roschin, A. V. (1967). "Toxicology of vanadium compounds used in modern industry". Gig Sanit. (Water Res.) 32: 26–32.
11
Vanadium [55] " Occupational Safety and Health Guidelines for Vanadium Pentoxide (http:/ / www. osha. gov/ SLTC/ healthguidelines/ vanadiumpentoxidedust/ recognition. html)". Occupational Safety and Health Administration. . Retrieved 2009-01-29. [56] Sax, N. I. (1984). Dangerous Properties of Industrial Materials, 6th ed.. Van Nostrand Reinhold Company. pp. 2717–2720. [57] N. B. Ress, B. J. Chou, R. A. Renne, J. A. Dill, R. A. Miller, J. H. Roycroft, J. R. Hailey, J. K. Haseman and J. R. Bucher (2003). "Carcinogenicity of Inhaled Vanadium Pentoxide in F344/N Rats and B6C3F1 Mice". Toxicological Sciences 74: 2876–296. doi: 10.1093/toxsci/kfg136 (http:/ / dx. doi. org/ 10. 1093/ toxsci/ kfg136). PMID 12773761. [58] Jörg M. Wörle-Knirsch, Katrin Kern, Carsten Schleh, Christel Adelhelm, Claus Feldmann, and Harald F. Krug (2007). "Nanoparticulate Vanadium Oxide Potentiated Vanadium Toxicity in Human Lung Cells". Environ. Sci. Technol. 41: 331–336. doi: 10.1021/es061140x (http:/ / dx. doi. org/ 10. 1021/ es061140x). [59] Ścibior, A.; Zaporowska, H.; Ostrowski, J. (2006). "Selected haematological and biochemical parameters of blood in rats after subchronic administration of vanadium and/or magnesium in drinking water". Archives of Environmental Contamination and Toxicology 51 (2): 287–295. doi: 10.1007/s00244-005-0126-4 (http:/ / dx. doi. org/ 10. 1007/ s00244-005-0126-4). [60] Gonzalez-Villalva, A.; et al. (2006). "Thrombocytosis induced in mice after subacute and subchronic V2O5 inhalation". Toxicology and Industrial Health 22 (3): 113–116. doi: 10.1191/0748233706th250oa (http:/ / dx. doi. org/ 10. 1191/ 0748233706th250oa). PMID 16716040. [61] Kazuo Kobayashia, Seiichiro Himeno, Masahiko Satoh, Junji Kuroda, Nobuo Shibata, Yoshiyuki Seko and Tatsuya Hasegawa (2006,). "Pentavalent vanadium induces hepatic metallothionein through interleukin-6-dependent and -independent mechanisms". Toxicology 228: 162–170. doi: 10.1016/j.tox.2006.08.022 (http:/ / dx. doi. org/ 10. 1016/ j. tox. 2006. 08. 022). [62] Soazo, Marina; Garcia, Graciela Beatriz (2007). "Vanadium exposure through lactation produces behavioral alterations and CNS myelin deficit in neonatal rats". Neurotoxicology and Teratology 29 (4): 503–510. doi: 10.1016/j.ntt.2007.03.001 (http:/ / dx. doi. org/ 10. 1016/ j. ntt. 2007. 03. 001). [63] Barceloux, Donald G.; Barceloux, Donald (1999). "Vanadium". Clinical Toxicology 37 (2): 265–278. doi: 10.1081/CLT-100102425 (http:/ / dx. doi. org/ 10. 1081/ CLT-100102425). [64] Ress, N. B.; et al. (2003). "Carcinogenicity of inhaled vanadium pentoxide in F344/N rats and B6C3F1 mice". Toxicological Sciences 74 (2): 287–296. doi: 10.1093/toxsci/kfg136 (http:/ / dx. doi. org/ 10. 1093/ toxsci/ kfg136). PMID 12773761. [65] Duffus, J. H. (2007). "Carcinogenicity classification of vanadium pentoxide and inorganic vanadium compounds, the NTP study of carcinogenicity of inhaled vanadium pentoxide, and vanadium chemistry". Regulatory Toxicology and Pharmacology 47 (1): 110–114. doi: 10.1016/j.yrtph.2006.08.006 (http:/ / dx. doi. org/ 10. 1016/ j. yrtph. 2006. 08. 006). [66] Opreskos, Dennis M. (1991). " Toxicity Summary for Vanadium (http:/ / rais. ornl. gov/ tox/ profiles/ old/ vanadium_f_V1. htm)". Oak Ridge National Laboratory. . Retrieved 2008-11-08. [67] http:/ / www. periodicvideos. com/ [68] http:/ / www. webelements. com/ webelements/ elements/ text/ V/ index. html [69] http:/ / www. atsdr. cdc. gov/ tfacts58. html
12
Article Sources and Contributors
Article Sources and Contributors Vanadium Source: http://en.wikipedia.org/w/index.php?oldid=306969096 Contributors: 1234321234bf, ATMyller, Ahoerstemeier, Alchemist-hp, Algebraist, Amgine, Andres, Apoc2400, Archimerged, Astaroth5, Atlant, Avenue, Axiosaurus, Baccyak4H, Beetstra, Benbest, Bender235, BerserkerBen, Bhowd, Biochemnick, BlueEarth, Bluezy, Borislav Dopudja, Breeden, Brian0918, Bryan Derksen, Bushytails, Cacahueten, Cadmium, Caeruleancentaur, Can't sleep, clown will eat me, Carnildo, CentaurWanderer, Chameleon, Charles.prescott, Chem-awb, Chrislk02, Conversion script, Coppertwig, Costelld, Crohnie, Crystal whacker, Darrien, David Latapie, Deafchild, Deglr6328, Delta G, DerHexer, Devilscrock, Discospinster, Donarreiskoffer, Dwmyers, Dyschunky, Eaolson, Eastlaw, Edgar181, Edgrmarriott, El, El aprendelenguas, Eldin raigmore, Emperorbma, Enigmaman, Fangfufu, Femto, Fivemack, Foobar, Frank Lofaro Jr., Frank Warmerdam, Gary King, Gekritzl, Gene Nygaard, Gh, Ginsengbomb, Glenn, Gman124, Gregogil, Grendelkhan, Gwarwick, Gwernol, Hak-kâ-ngìn, Hcefw, Hellbus, Herbee, Hut 8.5, Hydrogen Iodide, Icairns, Ideyal, Ingham, Ixfd64, J04n, Jaerik, Jakew, Janderk, Jaraalbe, Joanjoc, John, Jose77, Jqavins, Jsmaye, Kajasudhakarababu, Karl-Henner, Karlhahn, Kevenharlow, Koshogirl64, Kurykh, Kwamikagami, LA2, LAX, Lauren 218, Light current, Lizardo tx, Loren.wilton, LuigiManiac, Luk, Lumbercutter, MZMcBride, Malcolm Farmer, Mani1, Materialscientist, Matt Deres, Mav, McTrixie, Mccready, Mgimpel, Michael Devore, Milo99, Minesweeper, Mkweise, Mschel, Nergaal, Nihiltres, Oda Mari, Parslad, Pentalis, Petergans, Pftplzzz, Plazak, Poolkris, Poorjon, Psyche825, Pyrochem, RTC, Remember, Rich Farmbrough, Ridernyc, Rifleman 82, Rjhanson54, Rolinator, Romanm, Rror, Rursus, S3000, Sango123, Saperaud, Schneelocke, Sengkang, Sfuerst, Shafei, Shaneymcb, Shoeofdeath, Silverbackmarlin, Sl, Smokefoot, Snowmanradio, Sprklyuncrn, Squids and Chips, Srleffler, St loeffler, Stanleylhs, StaticGull, Stephenb, Stifynsemons, Stone, Strait, Suisui, Sunborn, Svante, Symplectic Map, Tagishsimon, Tarquin, Tetracube, The Voice Of Your Heart, Theseeker4, Thingg, Three-quarter-ten, Thricecube, Tiyoringo, Tomie16, Tony Fox, V1adis1av, V8rik, Vancouverguy, Veinor, Vlad4599, Vsmith, Walkerma, Warut, Wasell, Watch37264, Wiglaf, Wikipediarules2221, Xasodfuih, Yaf, Yarnalgo, Yersinia, Youkai no unmei, Yyy, Zaya05, Zinc2005, Zoicon5, 316 anonymous edits
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13
Chromium
1
Chromium vanadium ← chromium → manganese ↑ Cr ↓ Mo
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
24Cr Periodic table
Appearance silvery metallic
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
1656
1807
1991
2223
2530
2942
Chromium
2
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1590.6 kJ·mol−1 3rd: 2987 kJ·mol−1Atomic radiusCovalent radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusPoisson ratioMohs hardnessVickers hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of chromium iso
N.A.
half-life
DM
DE (MeV)
DP
50
4.345%
> 1.8×1017y
εε
-
50
51
syn
27.7025 d
ε
-
51
γ
0.320
-
Cr Cr
52
83.789%
52
53
9.501%
53
54
2.365%
54
Cr Cr Cr
Ti V
Cr is stable with 28 neutron Cr is stable with 29 neutron Cr is stable with 30 neutron
chromium, Cr, 24 transition metal6, 4, d51.9961(6) g·mol−1 [Ar] 3d5 4s1 2, 8, 13, 1 (Image) solid 7.19 g·cm−3 6.3 g·cm−3 2180 K,1907 °C,3465 °F 2944 K,2671 °C,4840 °F 21.0 kJ·mol−1 339.5 kJ·mol−1 (25 °C) 23.35 J·mol−1·K−16, 5, 4, 3, 2, 1, -1, -2 (strongly acidic oxide) 1.66 (Pauling scale) 1st: 652.9 kJ·mol−1128 pm139±5 pm body-centered cubic AFM (rather: SDW[1] ) (20 °C) 125 nΩ·m (300 K) 93.9 W·m−1·K−1 (25 °C) 4.9 µm·m−1·K−1 (20 °C) 5940 m/s 279 GPa 115 GPa 160 GPa 0.21 8.5 1060 MPa 1120 MPa 7440-47-3 Chromium (pronounced /ˈkroʊmiəm/) is a chemical element which has the symbol Cr and atomic number 24. It is a steely-gray, lustrous, hard metal that takes a high polish and has a high melting point. It is also odourless, tasteless, and malleable. The name of the element is derived from the Greek word "chrōma" (χρωμα), meaning color, because many of its compounds are intensely colored. It was discovered by Louis Nicolas Vauquelin in the mineral crocoite (lead chromate) in 1797. Crocoite was used as a pigment, and after the discovery that the mineral chromite also contains chromium this latter mineral was used to produce pigments as well. Chromium was regarded with great interest because of its high corrosion resistance and hardness. A major development was the discovery that steel could be made highly resistant to corrosion and discoloration by adding chromium and nickel to form stainless steel. This application, along with chrome plating (electroplating with chromium) are currently the highest-volume uses of the metal. Chromium and ferrochromium are produced from the single commercially viable ore, chromite, by silicothermic or aluminothermic reaction or by roasting and leaching processes. Although trivalent chromium (Cr(III)) is required in trace amounts for sugar and lipid metabolism in humans and its deficiency may cause a disease called chromium deficiency, hexavalent chromium (Cr(VI)) is toxic and carcinogenic, so that abandoned chromium production sites need environmental cleanup.
Chromium
3
Characteristics Occurrence Chromium is the 21st most abundant element in Earth's crust with an average concentration of 100 ppm.[2] Chromium compounds are found in the environment, due to erosion of chromium-containing rocks and can be distributed by volcanic eruptions. The concentrations range in soil is between 1 and 3000 mg/kg, in sea water 5 to 800 µg/liter, and in rivers and lakes 26 µg/liter to 5.2 mg/liter.[3] The relation between Cr(III) and Cr(VI) strongly depends on pH and oxidative properties of the location, but in most cases, the Cr(III) is the dominating species,[3] although in some areas the ground water can contain up to 39 µg of total chromium of which 30 µg is present as Cr(VI).[4] Chromium is mined as chromite (FeCr2O4) ore.[5] About two-fifths of the chromite ores and concentrates in the world are produced in South Africa, while Kazakhstan, India, Russia, and Turkey are also substantial producers. Untapped chromite deposits are plentiful, but geographically concentrated in Kazakhstan and southern Africa.[6] Though native chromium deposits are rare, some native chromium metal has been discovered.[7] [8] The Chromite ore Udachnaya Pipe in Russia produces samples of the native metal. This mine is a kimberlite pipe rich in diamonds, and the reducing environment helped produce both elemental chromium and diamond.[9]
Isotopes Naturally occurring chromium is composed of three stable isotopes; 52Cr, 53Cr and 54Cr with 52Cr being the most abundant (83.789% natural abundance). Nineteen radioisotopes have been characterized with the most stable being 50Cr with a half-life of (more than) 1.8x1017 years, and 51Cr with a half-life of 27.7 days. All of the remaining radioactive isotopes have half-lives that are less than 24 hours and the majority of these have half-lives that are less than 1 minute. This element also has 2 meta states.[10] 53
Cr is the radiogenic decay product of 53Mn. Chromium isotopic contents are typically combined with manganese isotopic contents and have found application in isotope geology. Mn-Cr isotope ratios reinforce the evidence from 26Al and 107Pd for the early history of the solar system. Variations in 53Cr/52Cr and Mn/Cr ratios from several meteorites indicate an initial 53Mn/55Mn ratio that suggests Mn-Cr isotopic composition must result from in-situ decay of 53Mn in differentiated planetary bodies. Hence 53Cr provides additional evidence for nucleosynthetic processes immediately before coalescence of the solar system.[11] The isotopes of chromium range in atomic mass from 43 u (43Cr) to 67 u (67Cr). The primary decay mode before the most abundant stable isotope, 52Cr, is electron capture and the primary mode after is beta decay.[10]
Chromium
4
Chemistry Oxidation states [12] [13] of chromium −2
Na2[Cr(CO)5]
−1
Na2[Cr2(CO)10]
0
Cr(C6H6)2
+1
K3[Cr(CN)5NO]
+2
CrCl2
+3
CrCl3
+4
K2CrF6
+5
K3CrO8
+6
K2CrO4
Chromium is a member of the transition metals, in group 6. Chromium(0) has an electronic configuration of 4s13d5, due to the lower energy of the high spin configuration. Chromium exhibits a wide range of possible oxidation states. The most common oxidation states of chromium are +2, +3, and +6, with +3 being the most stable. +1, +4 and +5 are rare. Chromium compounds of oxidation state +6 are powerful oxidants. All (except the hexafluoride and chromium hexacarbonyl) stable chromium compounds of the oxidation state +6 contain oxygen as ligand, for example the chromate (CrO42-) and Chromyl chloride (CrO2Cl2).[15]
Chemistry and compounds The oxidation state 3+ is the most stable one and therefore a large number of The Pourbaix diagram for chromium in pure water, chromium(III) compounds is known. [3] [14] perchloric acid or sodium hydroxide Chromium(III) can be obtained by dissolving chromium in acids like hydrochloric acid or sulfuric acid. The aluminium(III) (ion radius 0.50 Å) and chromium(III) (ion radius 0.63 Å) can replace each other in some compounds, for example chrome alum and alum. Another example is aluminium oxide (corundum, Al2O3) where by replacement, the red colored ruby is formed.
Chromium
5
Chromium tends to form complexes, for example with water molecules (hydrates); the chromium ions in water are usually octahedrally coordinated. The commercially available chromium(III) chloride hydrate is the dark green complex [CrCl2(H2O)4]Cl, but two other forms are known, viz., pale green [CrCl(H2O)5]Cl2 and violet [Cr(H2O)6]Cl3. If water-free green chromium(III) chloride is dissolved in water then the green solution turn violet after some time. This color change is due to the substitution of water for chloride in the inner coordination sphere. This kind of reactions is also Chromium(III) chloride hexahydrate observed in chrome alum solutions and some other ([CrCl2(H2O)4]Cl•2H2O) water soluble chromium(III) salts, and the reverse reaction can be induced by heating the solution. The chromium(III) hydroxide Cr(OH)3 shows amphoteric reactions and dissolves in acid water by forming [Cr(H2O)6]3+ and in basic water by forming [Cr(OH)6]3-. By heating the chromium(III) hydroxide it is transformed into the green chromium(III) oxide (Cr2O3), which is the stable oxide (melting point of 2275 °C) with the crystal structure identical to that of corundum.[15] The second stable oxidation state is 6+, for example the chromate, which is produced in large scale by oxidative roasting of chromite ore with calcium or sodium carbonate. Chromate and dichromate are in an equilibrium, which is influenced in that case by the law of mass action and therefore by the pH of the solution. 2-
2 CrO
4
+ 2 H O+ → Cr O 3
2
27
+2H O 2
The change in equilibrium is also visible by a change
Chromium(III) chloride (CrCl3)
from yellow (chromate) to orange (dichromate) if an acid is added to a neutral solution of potassium chromate. At lower pH, further condensation to more complex oxyanions of chromium is possible. The chromate and dichromate are strong oxidizing reagents at low pH[15] Cr2O72- + 14 H3O+ + 6 e- → 2 Cr3+ + 21 H2O ε0 = 1.33V but only moderate ones at high pH[15] CrO42- + 4 H2O + 3 e- → Cr(OH)3+ + 5 OH- ε0 = - 0.13V.
Chromium(VI) oxide
Chromium(VI) compounds in solution can be detected by adding acidic hydrogen peroxide solution. A dark blue unstable chromium(VI) peroxide (CrO5) is formed which can be stabilized as an ether adduct CrO5 • OR2.[15] Chromic acid has the hypothetical structure H2CrO4. Neither chromic nor dichromic acid is found in nature, but their anions are found in a variety of compounds, the chromates and dichromates. chromium(VI) oxide CrO3, the acid anhydride of
Chromium chromic acid, is sold industrially as "chromic acid".[15] The dark red chromium(VI) oxide can be produced by mixing sulfuric acid with dichromate, and is an extremely strong oxidizing agent. The oxidation state 5+ is only realized in few compounds. The only binary compound is the highly volatile chromium(V) fluoride (CrF5). This red solid with a melting point of 30°C and a boiling point of 117°C can be synthesized by reacting fluorine with chromium at 400°C and 200 bar pressure. The peroxochromate(V) is another example of the oxidation state 5+. The potassium peroxochromate (K3[Cr(O2)4]) is made by Sodium chromate reacting potassium chromate with hydrogen peroxide at low temperatures. This red brown compound is stable at room temperature but decomposes spontaneously at 150–170 °C.[16] The chromium(IV) compounds (4+) are slightly more stable than the chromium(V) compounds, and the halogen compounds CrF4, CrCl4 and CrBr4 can be produced by the reaction of the trihalogens with additional elementary halogens at elevated temperatures. Most of the compounds are susceptible to disproportionation reactions and therefore are not stable in water. An example for a Chromium(II) compounds (2+) is the water stable chromium(II) chloride which can be produced by reduction of chromium(III) chloride with zinc. The resulting light blue solutions are only stable at neutral pH when the solution is very pure.[15] Passivation Chromium is passivated by oxygen, forming a thin protective oxide surface layer. This layer is a spinel structure only a few atoms thick. It is very dense, preventing diffusion of oxygen into the underlying material. (In iron or plain carbon steels the oxygen migrates into the underlying material.)[17] Chromium is usually plated on top of a nickel layer which may first have been copper plated.[15] Chromium, unlike metals such as iron and nickel, does not suffer from hydrogen embrittlement. It does suffer from nitrogen embrittlement - chromium reacts with nitrogen from air and forms brittle nitrides at temperatures necessary to work the metal parts [18] . The Pourbaix diagram can be seen above. It is important to understand that the diagram only displays the thermodynamic data and it does not display any details of the rates of reaction.[3] The passivation can be increased by short contact with oxidizing acids like nitric acid. The passivated chromium is stable against acids. The contrary effect can be achieved if a strong reducing reactant destroys the oxide protection layer on the metal, a metal treated in this way readily dissolves in weak acids.[15]
6
Chromium
7
Quintuple bond Chromium is notable for its ability to form quintuple covalent bonds. The synthesis of a compound of chromium(I) and a hydrocarbon radical was shown via X-ray diffraction to contain a quintuple bond of length 183.51(4) pm (1.835 angstroms) joining the two central chromium atoms.[19] Extremely bulky monodentate ligands prevent that bonds to other atoms can be formed and therefore stabilizes this compound.
Physical properties Chromium is remarkable for its magnetic properties: it is the only elemental solid which shows antiferromagnetic ordering at room temperature (and below). Above 38 °C, it transforms into a paramagnetic state [1] .
Chromium compound, which was determined experimentally to contain a Cr-Cr quintuple bond
History Weapons found in burial pits dating from the late 3rd century BC Qin Dynasty of the Terracotta Army near Xi'an, China have been analyzed by archaeologists. Although buried more than 2,000 years ago, the ancient bronze tips of crossbow bolts and swords found at the site showed no sign of corrosion, because the bronze was coated with chromium.[20] Chromium came to the attention of westerners in the
Crocoite (PbCrO4)
18th century. On 26 July 1761, Johann Gottlob Lehmann found an orange-red mineral in the Beryozovskoye mines in the Ural Mountains which he named Siberian red lead. Though misidentified as a lead compound with selenium and iron components, the mineral was Crocoite (lead chromate) with a formula of PbCrO4.[21] In 1770, Peter Simon Pallas visited the same site as Lehmann and found a red lead mineral that had useful properties as a pigment in paints. The use of Siberian red lead as a paint pigment developed rapidly. A bright yellow pigment made from crocoite also became fashionable.[21]
Chromium
8 In 1797, Louis Nicolas Vauquelin received samples of crocoite ore. He produced chromium oxide (CrO3) by mixing crocoite with hydrochloric acid. In 1798, Vauquelin discovered that he could isolate metallic chromium by heating the oxide in a charcoal oven.[22] He was also able to detect traces of chromium in precious gemstones, such as ruby or emerald.[21] [23]
During the 1800s, chromium was primarily used as a component of paints and in tanning salts. At first, crocoite from Russia was the main source, but in 1827, Ruby is colored by a small amount of a larger chromite deposit was discovered near chromium Baltimore, United States. This made the United states the largest producer of chromium products till 1848 when large deposits of chromite where found near Bursa, Turkey.[5] Chromium is also known for its luster when polished. It is used as a protective and decorative coating on car parts, plumbing fixtures, furniture parts and many other items, usually applied by electroplating. Chromium was used for electroplating as early as 1848, but this use only became widespread with the development of an improved process in 1924.[24] Metal alloys now account for 85% of the use of chromium. The remainder is used in the chemical industry and refractory and foundry industries.
Production Approximately 4.4 million metric tons of marketable chromite ore were produced in 2000, and converted into ~3.3 million tons of ferro-chrome with an approximate market value of 2.5 billion United States dollars.[25] The largest producers of chromium ore have been South Africa (44%) India (18%), Kazakhstan (16%) Zimbabwe (5%), Finland (4%) Iran (4%) and Brazil (2%) with several other countries producing the rest of less than 10% of the world production.[25]
World production trend of chromium
The two main products of chromium ore refining are ferrochromium and metallic chromium. For those products the ore smelter process differs considerably. For the production of ferrochromium, the chromite ore (FeCr2O4) is reduced in large scale in electric arc furnace or in smaller smelters with either aluminium or silicon in an aluminothermic reaction.[26]
Chromium
For the production of pure chromium, the iron has to be separated from the chromium in a two step roasting and leaching process. The chromite ore is heated with a mixture of calcium carbonate and sodium carbonate in the presence of air. The chromium is oxidized to the [25] Chromium ore output in 2002 hexavalent form, while the iron forms the stable Fe2O3. The subsequent leaching at higher elevated temperatures dissolves the chromates and leaves the insoluble iron oxide. The chromate is converted by sulfuric acid into the dichromate.[26] 4FeCr2O4 + 8 Na2CO3 + 7 O2 → 8 Na2CrO4 + 2 Fe2O3 + 8 CO2 2Na2CrO4 + H2SO4 → Na2Cr2O7 + Na2SO4 + H2O The dichromate is converted to the chromium(III) oxide by reduction with carbon and then reduced in an aluminothermic reaction to chromium.[26] Na2Cr2O7 + 2 C → Cr2O3 + Na2CO3 + CO Cr2O3 + 2 Al → Al2O3 + Cr
Applications Metallurgy The strengthening effect on steel by forming stable carbide grains at the grain boundaries and the strong increase in corrosion resistance made chromium an important alloying material for steel. The high speed tool steels contain between 3 and 5% chromium. An important stainless steel is 18/10 stainless, made from iron with 10% nickel and 18% chromium, is widely used for cookware and cutlery. For these applications, ferrochromium is added to the molten iron. Also Decorative chrome plating on a nickel-based alloys increase in strength due to the motorcycle. formation of stable carbide grains at the grain boundaries. For example, Inconel 718 contains 18.6% chromium. Because of the excellent heat stability of these nickel superalloys, they are used in jet engines and gas turbines in large quantities.[27] The relative high hardness and corrosion resistance of unalloyed chromium makes it a good surface coating. A thin layer of chromium is deposited on pretreated metallic surfaces by electroplating techniques. There are two deposition methods: Thin, below 1 µm thickness, layers are deposited by chrome plating, and are used for decorative surfaces. If wear-resistant surfaces are needed then thicker chromium layers of up to mm thickness are deposited. Both methods normally use acidic chromate or dichromate solutions. To prevent the energy consuming change in oxidation state, the use of Chromium(III) sulfate is under development, but for most applications, the established process is used.[24]
9
Chromium
10
In the chromate conversion coating process, the strong oxidative properties of chromates are used to deposit a protective oxide layer on metals like aluminium, zinc and cadmium. This passivation and the self healing properties by the chromate stored in the chromate conversion coating, which is capable to migrate to local defects, are the benefits of this [28] coating method. Because of environmental and health regulations on chromates, alternative coating method are under development.[29] Anodizing of aluminium is another electrochemical process, which does not lead to the deposition of chromium, but uses chromic acid as electrolyte in the solution. During anodization, an oxide layer is formed on the aluminium. The use of chromic acid, instead of the normally used sulfuric acid, leads to a slight difference of these oxide layers.[30] The high toxicity of Cr(VI) compounds, used in the established chromium electroplating process, and the strengthening of safety and environmental regulations demand a search for substitutes for chromium or at least a change to less toxic chromium(III) compounds.[24]
Dye and pigment The mineral crocoite (lead chromate PbCrO ) was used 4
as a yellow pigment shortly after its discovery. After a synthesis method became available starting from the more abundant chromite, Chrome yellow was, together with cadmium yellow, one of the most used yellow pigments. The pigment does not degrade in the School bus painted in Chrome [31] yellow light and has a strong color. The signaling effect of yellow was used for school buses in the United States and for Postal Service (for example Deutsche Post) in Europe. The use of chrome yellow declined due to environmental and safety concerns and was substituted by organic pigments or other lead-free alternatives.[32] Other pigments based on chromium are, for example, the bright red pigment Chrome red, which is a basic lead chromate (PbCrO4•Pb(OH)2).[32] Chrome green is a mixture of Prussian blue and chrome yellow, while the Chrome oxide green is Chromium(III) oxide.[32] Glass is colored green by the addition of chromium(III) oxide. This is similar to emerald, which is also colored by chromium.[33] A red color is achieved by doping chromium(III) into the crystals of corundum, which are then called ruby. Therefore, chromium is used in producing synthetic rubies.[34] The toxicity of chromium(VI) salts is used in the preservation of wood. For example, chromated copper arsenate (CCA) is used in timber treatment to prevent wood from decay fungi, wood attacking insects, including termites, and marine borers.[35] The formulations contain chromium based on the oxide CrO3 between 35.3% and 65.5%. In the United States, 65,300 metric tons of CCA solution have been used in 1996.[35]
Chromium
Tanning Chromium(III) salts, especially chrome alum and chromium(III) sulfate, are used in the tanning of leather. The chromium(III) stabilizes the leather by cross linking the collagen fibers within the leather.[36] Chromium tanned leather can contains between 4 and 5% of chromium, which is tightly bound to the proteins.[5]
Refractory material The high heat resistivity and high melting point makes chromite and chromium(III) oxide a material for high temperature refractory applications, like blast furnaces, cement kilns, molds for the firing of bricks and as foundry sands for the casting of metals. In these applications, the refractory materials are made from mixtures of chromite and magnesite. The use is declining because of the environmental regulations due to the possibility of the formation of chromium(VI).[26]
Other use Several chromium compounds are used as catalyst. For example the Phillips catalysts for the production polyethylene are mixtures of chromium and silicon dioxide or mixtures of chromium and titanium and aluminium oxide.[37] Chromium(IV) oxide (CrO2) is a magnetic compound. Its ideal shape anisotropy, which imparted high coercivity and remanent magnetization, made it a compound superior to the γ-Fe2O3. Chromium(IV) oxide is used to manufacture magnetic tape used in high performance audio tape and standard audio cassette.[38] Chromates can prevent corrosion of steel under wet conditions, and therefore chromates are added to the drilling muds.[39] The long known influence of chromium uptake on diabetes conditions suggested the positive influence of dietary supplement containing chromium(III) also on healthy persons. For this reason, dietary supplement or slimming aid usually contain chromium(III) chloride, chromium(III) picolinate, chromium(III) polynicotinate or amino acid chelate, such as chromium(III) D-phenylalanine. The benefit of those supplements is still under investigation and is questioned by some studies.[40] [41] • Chromium hexacarbonyl Cr(CO)6 is used as a gasoline additive.[42] • Chromium(III) oxide is a metal polish known as green rouge. • Chromic acid is a powerful oxidizing agent and is a useful compound for cleaning laboratory glassware of any trace of organic compounds. It is prepared in situ by dissolving potassium dichromate in concentrated sulfuric acid, which is then used to wash the apparatus. Sodium dichromate is sometimes used because of its higher solubility (5 g/100 ml vs. 20 g/100 ml respectively). Potassium dichromate is a chemical reagent, used in cleaning laboratory glassware and as a titrating agent. It is also used as a mordant (i.e., a fixing agent) for dyes in fabric.
Biological role Trivalent chromium (Cr(III) or Cr3+) is required in trace amounts for sugar and lipid metabolism in humans, and its deficiency may cause a disease called chromium deficiency.[43] In contrast, hexavalent chromium (Cr(VI) or Cr6+) is very toxic and mutagenic when inhaled. Cr(VI) has not been established as a carcinogen when in solution, though it may cause allergic contact dermatitis (ACD).[44]
11
Chromium The use of chromium-containing dietary supplements is controversial due to the complex effects of the used supplements.[45] The popular dietary supplement chromium picolinate complex generates chromosome damage in hamster cells.[46] In the United States the dietary guidelines for daily chromium uptake were lowered from 50-200 µg for an adult to [47] 35 µg (adult male) and to 25 µg (adult female).
Precautions Water insoluble chromium(III) compounds and chromium metal are not considered a health hazard, while the toxicity and carcinogenic properties of chromium(VI) are known for a long time.[48] An actual investigation into hexavalent chromium release into drinking water was used as the plot-basis of the motion picture Erin Brockovich. Because of the specific transport mechanisms, only limited amounts of chromium(III) enter the cells. Several in vitro studies indicated that high concentrations of chromium(III) in the cell can lead to DNA damage.[49] Acute oral toxicity ranges between 1500 and 3300 µg/kg.[50] The proposed beneficial effects of chromium(III) and the use as dietary supplements yielded some controversial results, but recent reviews suggest that moderate uptake of chromium(III) through dietary supplements poses no risk.[49] World Health Organization recommended maximum allowable concentration in drinking water for chromium (VI) is 0.05 milligrams per liter. Hexavalent chromium is also one of the substances whose use is restricted by the European Restriction of Hazardous Substances Directive. The acute oral toxicity for chromium(VI) ranges between 50 and 150 µg/kg.[50] In the body, chromium(VI) is reduced by several mechanisms to chromium(III) already in the blood before it enters the cells. The chromium(III) is excreted from the body, whereas the chromate ion is transferred into the cell by a transport mechanism, by which also sulfate and phosphate ions enter the cell. The acute toxicity of chromium(VI) is due to its strong oxidational properties. After it reaches the blood stream, it damages the kidneys, the liver and blood cells through oxidation reactions. Hemolysis, renal and liver failure are the results of these damages. Aggressive dialysis can improve the situation.[51] The carcinogenity of chromate dust is known for a long time, and in 1890 the first publication described the elevated cancer risk of workers in a chromate dye company.[52] [53] Three mechanisms have been proposed to describe the genotoxicity of chromium(VI). The first mechanism includes highly reactive hydroxyl radicals and other reactive radicals which are by products of the reduction of chromium(VI) to chromium(III). The second process includes the direct binding of chromium(V), produced by reduction in the cell, and chromium(IV) compounds to the DNA. The last mechanism attributed the genotoxicity to the binding to the DNA of the end product of the chromium(III) reduction.[54] Chromium salts (chromates) are also the cause of allergic reactions in some people. Chromates are often used to manufacture, amongst other things, leather products, paints, cement, mortar and anti-corrosives. Contact with products containing chromates can lead to allergic contact dermatitis and irritant dermatitis, resulting in ulceration of the skin, sometimes referred to as "chrome ulcers". This condition is often found in workers that have been exposed to strong chromate solutions in electroplating, tanning and chrome-producing manufacturers.[55] [56] [56] In some parts of Russia, pentavalent chromium was reported as one of the causes of premature senility.[57]
12
Chromium
Environmental issues As chromium compounds were used in dyes and paints and the tanning of leather, these compounds are often found in soil and groundwater at abandoned industrial sites, now needing environmental cleanup and remediation per the treatment of brownfield land. Primer paint containing hexavalent chromium is still widely used for aerospace and automobile refinishing applications.
See also • • • •
Chromium compounds Chromium minerals Chromium(III) picolinate Chromium VI
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Chromium Research Council (U.S.). Committee on Coatings, Published by National Academy of Sciences, 1970 ISBN 0309017696, 9780309017695 [19] T. Nguyen, A. D. Sutton, M. Brynda, J. C. Fettinger, G. J. Long and P. P. Power (2005). "Synthesis of a Stable Compound with Fivefold Bonding Between Two Chromium(I) Centers". Science 310 (5749): 844–847. doi: 10.1126/science.1116789 (http:/ / dx. doi. org/ 10. 1126/ science. 1116789). PMID 16179432. [20] Cotterell, Maurice. (2004). The Terracotta Warriors: The Secret Codes of the Emperor's Army. Rochester: Bear and Company. ISBN 159143033X. Page 102. [21] Jacques Guertin, James Alan Jacobs, Cynthia P. Avakian, (2005). Chromium (VI) Handbook. CRC Press. pp. 7–11. ISBN 9781566706087. [22] Vauquelin, Louis Nicolas (1798). " Memoir on a New Metallic Acid which exists in the Red Lead of Sibiria (http:/ / books. google. com/ books?id=6dgPAAAAQAAJ)". Journal of Natural Philosophy, Chemistry, and the Art 3: 146. . [23] van der Krogt, Peter, Chromium (http:/ / elements. vanderkrogt. net/ elem/ cr. html), , retrieved 2008-08-24 [24] Dennis, J. K.; Such, T. E. (1993). "History of Chromium Plating". Nickel and Chromium Plating. Woodhead Publishing. pp. 9–12. ISBN 9781855730816. [25] Papp, John F. " Mineral Yearbook 2002: Chromium (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ chromium/ chrommyb02. pdf)". United States Geological Survey. . Retrieved 2009-02-16. [26] Papp, John F; Lipin Bruce R. (2006). " Chromite (http:/ / books. google. de/ books?id=zNicdkuulE4C& pg=PA309)". Industrial Minerals & Rocks: Commodities, Markets, and Uses (7th ed.). SME. ISBN 9780873352338. . [27] Bhadeshia, H. K. D. H.. " Nickel-Based Superalloys (http:/ / www. msm. cam. ac. uk/ phase-trans/ 2003/ Superalloys/ superalloys. html)". University of Cambridge. . Retrieved 2009-02-17. [28] Edwards, Joseph (1997). Coating and Surface Treatment Systems for Metals. Finishing Publications Ltd. and ASM International. pp. 66–71. ISBN 0-904477-16-9. [29] Zhao, J.; Xia, L.; Sehgal, A.; Lu, D.; McCreery, R. L.; Frankel, G. S. (2001). " Effects of chromate and chromate conversion coatings on corrosion of aluminum alloy 2024-T3 (http:/ / www. chemistry. ohio-state. edu/ ~rmccreer/ group/ mccreery157. pdf)". Surface and Coatings Technology 140 (1): 51–57. doi: 10.1016/S0257-8972(01)01003-9 (http:/ / dx. doi. org/ 10. 1016/ S0257-8972(01)01003-9). . [30] Sprague, J. A.; Smidt, F. A. (1994). ASM Handbook: Surface Engineering (http:/ / books. google. com/ books?id=RGtsPjqUwy0C& pg=PA484). ASM International. ISBN 9780871703842. . Retrieved 2009-02-17. [31] Worobec, Mary Devine; Hogue, Cheryl (1992). Toxic Substances Controls Guide: Federal Regulation of Chemicals in the Environment (http:/ / books. google. de/ books?id=CjWQ6_7AnI4C& pg=PA13). Washington, D.C.: Bureau of National Affairs. p. 13. ISBN 9780871797520. . [32] Gettens, Rutherford John; Stout, George Leslie (1966). Painting Materials: A Short Encyclopaedia (http:/ / books. google. de/ books?id=bdQVgKWl3f4C& pg=PA106). Courier Dover Publications. pp. 105-–106. ISBN 9780486215976. . [33] Carstens, Harald (1973). "The red-green change in chromium-bearing garnets". Contributions to Mineralogy and Petrology 41 (3): 273–276. doi: 10.1007/BF00371036 (http:/ / dx. doi. org/ 10. 1007/ BF00371036). [34] Moss, S. C.; Newham, R. E. (1964). " The chromium position in ruby (http:/ / rruff. geo. arizona. edu/ doclib/ zk/ vol120/ ZK120_359. pdf)". Zeitschrift fur Kristallographie 120: 359–363. . [35] "Leaching of chromated copper arsenate wood preservatives: a review". Environmental Pollution 111 (1): 53–66. 2001. doi: 10.1016/S0269-7491(00)00030-0 (http:/ / dx. doi. org/ 10. 1016/ S0269-7491(00)00030-0). [36] Brown, E. M.; Dudley, R.L.; Elsetinow A. R. (1997). "A Conformational Study of Collagen as Affected by Tanning Procedures". Journal of the American Leather Chemists Association 92: 225–233. [37] Weckhuysen, Bert M.; Schoonheydt Robert A. (1999). "Olefin polymerization over supported chromium oxide catalysts". Catalysis Today 51 (2): 215–221. doi: 10.1016/S0920-5861(99)00046-2 (http:/ / dx. doi. org/ 10. 1016/ S0920-5861(99)00046-2). [38] Mallinson, John C. (1993). " Chromium Dioxide (http:/ / books. google. de/ books?id=rNifWsBxnWkC& pg=PA32)". The foundations of magnetic recording. Academic Press. ISBN 9780124666269. . [39] Garverick, Linda (1994). Corrosion in the Petrochemical Industry (http:/ / books. google. de/ books?id=qTfNZZRO758C& pg=PA278). ASM International. ISBN 9780871705051. . [40] Heimbach, J.T.; Anderson, R.A. (2005). " Chromium: Recent Studies Regarding Nutritional Roles and Safety (http:/ / journals. lww. com/ nutritiontodayonline/ Abstract/ 2005/ 07000/ Chromium__Recent_Studies_Regarding_Nutritional. 13. aspx)". Nutrition Today 40 (4): 189–-195. . [41] Vincent,, John B . (2003). "The Potential Value and Toxicity of Chromium Picolinate as a Nutritional Supplement, Weight Loss Agent and Muscle Development Agent". Sports Medicine:Volume 33 (3): 213–230. doi: 10.2165/00007256-200333030-00004 (http:/ / dx. doi. org/ 10. 2165/ 00007256-200333030-00004). [42] Patnaik, Pradyot (2003). " Chromium hexacarbonyl (http:/ / books. google. de/ books?id=Xqj-TTzkvTEC& pg=PA222)". Handbook of Inorganic Chemicals. McGraw-Hill Professional. pp. 222–223. ISBN 9780070494398.
14
Chromium . [43] Mertz, Walter (01 April 1993). " Chromium in Human Nutrition: A Review (http:/ / jn. nutrition. org/ cgi/ content/ abstract/ 123/ 4/ 626)". Journal of Nutrition 123 (4): 626–636. PMID 8463863. . [44] " ToxFAQs: Chromium (http:/ / www. atsdr. cdc. gov/ tfacts7. html)". Agency for Toxic Substances & Disease Registry, Centers for Disease Control and Prevention. February 2001. . Retrieved 2007-10-02. [45] Cronin, Joseph R. (2004). "The Chromium Controversy". Alternative and Complementary Therapies 10 (1): 39–42. doi: 10.1089/107628004772830393 (http:/ / dx. doi. org/ 10. 1089/ 107628004772830393). [46] Stearns, D. M.; Wise, J. P.; Patierno, S. R.; Wetterhahn, K. E. (01 December 1995). " Chromium(III) picolinate produces chromosome damage in Chinese hamster ovary cells (http:/ / www. fasebj. org/ cgi/ content/ abstract/ 9/ 15/ 1643)". Federation of American Societies for Experimental Biology 9 (15): 1643–1648. PMID 8529845. . [47] Vincent, J. B. (2007). "Recent advances in the nutritional biochemistry of trivalent chromium". Proceedings of the Nutrition Society 63 (01): 41–47. doi: 10.1079/PNS2003315 (http:/ / dx. doi. org/ 10. 1079/ PNS2003315). [48] Barceloux, Donald G.; Barceloux, Donald (1999). "Chromium". Clinical Toxicology 37 (2): 173–194. doi: 10.1081/CLT-100102418 (http:/ / dx. doi. org/ 10. 1081/ CLT-100102418). [49] Eastmond, David A.; MacGregor, James T.; Slesinski, Ronald S. (2008). "Trivalent Chromium: Assessing the Genotoxic Risk of an Essential Trace Element and Widely Used Human and Animal Nutritional Supplement". Critical Reviews in Toxicology 38 (3): 173–190. doi: 10.1080/10408440701845401 (http:/ / dx. doi. org/ 10. 1080/ 10408440701845401). [50] Katz, Sidney A.; Salem Harry (1992). "The toxicology of chromium with respect to its chemical speciation: A review". Journal of Applied Toxicology 13 (3): 217–224. doi: 10.1002/jat.2550130314 (http:/ / dx. doi. org/ 10. 1002/ jat. 2550130314). [51] Dayan, A. D.; Paine, A. J. (2001). "Mechanisms of chromium toxicity, carcinogenicity and allergenicity: Review of the literature from 1985 to 2000". Human & Experimental Toxicology 20 (9): 439–451. doi: 10.1191/096032701682693062 (http:/ / dx. doi. org/ 10. 1191/ 096032701682693062). PMID 11776406. [52] Newman, D. (1890). "A case of adeno-carcinoma of the left inferior turbinated body, and perforation of the nasal septum, in the person of a worker in chrome pigments". Glasgow Med J 33: 469–470. [53] Langard, Sverre (1990). "One Hundred Years of Chromium and Cancer: A Review of Epidemiological Evidence and Selected Case Reports". American Journal of Industrial Medicine 17: 189–215. doi: 10.1002/ajim.4700170205 (http:/ / dx. doi. org/ 10. 1002/ ajim. 4700170205). [54] M. D., Cohen; Kargacin, B.; Klein, C. B.; Costa, M. (1993). "Mechanisms of chromium carcinogenicity and toxicity". Critical reviews in toxicology 23 (3): 255–81. doi: 10.3109/10408449309105012 (http:/ / dx. doi. org/ 10. 3109/ 10408449309105012). [55] " Chrome Contact Allergy (http:/ / dermnetnz. org/ dermatitis/ chrome-allergy. html)". DermNet NZ. . [56] Basketter, David; Horev, Liran; Slodovnik, Dany; Merimes, Sharon; Trattner, Akiva; Ingber, Arieh (2000). "Investigation of the threshold for allergic reactivity to chromium". Contact Dermatitis 44 (2): 70–74. doi: 10.1034/j.1600-0536.2001.440202.x (http:/ / dx. doi. org/ 10. 1034/ j. 1600-0536. 2001. 440202. x). [57] Chromium Toxicity (http:/ / www. corrosion-doctors. org/ Pollution/ chromiumtoxicity. htm) on the Corrosion Doctors Web site maintained by Canadian Physical Chemist, Pierre R. Roberge, PhD, P.Eng. (access date 27 april 2009)
External links • ATSDR Case Studies in Environmental Medicine: Chromium Toxicity (http:/ / www. atsdr. cdc. gov/ csem/ chromium) U.S. Department of Health and Human Services • Los Alamos National Laboratory - Chromium (http:/ / periodic. lanl. gov/ elements/ 24. html) • WebElements.com – Chromium (http:/ / www. webelements. com/ webelements/ elements/ text/ Cr/ index. html) • IARC Monograph "Chromium and Chromium compounds" (http:/ / www-cie. iarc. fr/ htdocs/ monographs/ vol49/ chromium. html) • International Chromium Development Association (http:/ / www. chromium-asoc. com/ ) • It's Elemental – The Element Chromium (http:/ / education. jlab. org/ itselemental/ ele024. html) • National Pollutant Inventory - Chromium (III) compounds fact sheet (http:/ / www. npi. gov. au/ database/ substance-info/ profiles/ 24. html)
15
Chromium • The Merck Manual – Mineral Deficiency and Toxicity (http:/ / www. merck. com/ mmpe/ sec01/ ch005/ ch005b. html) • National Institute for Occupational Safety and Health - Chromium Page (http:/ / www. cdc. gov/ niosh/ topics/ chromium/ )
16
Article Sources and Contributors
Article Sources and Contributors Chromium Source: http://en.wikipedia.org/w/index.php?oldid=307638934 Contributors: 12dstring, A2Kafir, AHEMSLTD, Accurizer, Adashiel, Addshore, Ahoerstemeier, Aitias, Akradecki, Ale jrb, Alecmconroy, Alex43223, Alexcalamaro,
[email protected], AndonicO, Andres, Antandrus, Anwar saadat, Apparition11, Arcadian, Archimerged, Artichoker, Athaler, Atjesse, Audrius u, Aussie Alchemist, Bbernet13, Bcorr, Beetstra, BenFrantzDale, Benbest, Bently34, BillFlis, Billwiki2008, Binary TSO, Blastwizard, Bobo192, Bogey97, Bomac, Bork, Bowlhover, Brian Huffman, Brian0918, Bryan Derksen, Burritobeatle, Cacycle, Cadmium, Can't sleep, clown will eat me, CanadianLinuxUser, Capricorn42, CaptainP, Carnildo, Cbc21com, Cenaboy2054, Chameleon, ChemNerd, Chiu frederick, ChrisnHouston, Chrissy385, Cometstyles, Conversion script, Coppertwig, CorvetteZ51, Cremepuff222, Curps, Danelo, DanielCristofani, Dantelebeau, Daqu, Darrien, Davewild, David Latapie, DavidCary, Davidmello, Davidruben, Decatoncale, Deli nk, Delldot, Delta G, DerHexer, Deville, Dhp1080, Discospinster, Donarreiskoffer, DoubleBlue, Doulos Christos, Drhaggis, Duk, Dwmyers, E Wing, Edgar181, Ekke44, El C, Eldin raigmore, Elkman, Elvey, Emperorbma, Epbr123, EstebanF, Euryalus, Everyking, Femto, Fifo, Frank Warmerdam, Fratrep, Fredrik, FrenchIsAwesome, Frotz, Fuosing, GRAHAMUK, Gaius Cornelius, Gene Nygaard, Geno-Supremo, Giftlite, Gilliam, Gimboid13, Gman124, Gordjazz, Greatpatton, Grendelkhan, Gurch, Gwernol, H Bruthzoo, Haddendaddendoedenda, Hak-kâ-ngìn, HappyM, Haza-w, Hellbus, Heron, Hobartimus, Hondai, Hroðulf, Hurmata, Hut 6.5, Hut 8.5, Hydrogen Iodide, Ian Pitchford, Icairns, Iridescent, Isindil, Ixfd64, J.delanoy, Jake Wartenberg, James086, Jaraalbe, Jaxl, Jj137, Joanjoc, John, John Bahrain, John Doe or Jane Doe, Jose77, Jsonitsac, Karl-Henner, Karlhahn, Keilana, Kjkolb, Knife Knut, Krawi, Kurykh, Kwamikagami, Kyanite, LA2, La goutte de pluie, Lankiveil, Larrybobb, Leebo, Lethalgeek, Leyo, Lightblade, LinguisticDemographer, Little Mountain 5, Lordryker, LorenzoB, Luigi30, LuigiManiac, Luwilt, MPerel, MZMcBride, Mackant1, Madlobster, MarkRose, Markhurd, Materialscientist, Mav, Meeples, Mgimpel, Michael Devore, Michaelbusch, Midnight Comet, Mikelieman, Mikeo, Minesweeper, Mm40, NRen2k5, Naffer, NawlinWiki, Nergaal, Nick Y., NickelShoe, Nihiltres, Novel tubes, NuclearWarfare, OccamzRazor, Omegatron, Omicronpersei8, OrbitOne, Otisjimmy1, Oxymoron83, PericlesofAthens, Persian Poet Gal, Pgk, Philip Trueman, Physchim62, Plexust, Poeloq, Poolkris, Potterfa11, PrestonH, Pretzelpaws, Prodego, Pseudopanax, Psyche825, Pzavon, Quadrius, Quintote, R'n'B, RJaguar3, RTC, Rada, Rallette, Ravichandar84, Red Alien, Remember, Renaissancee, RexNL, Rfc1394, Rich Farmbrough, Richard D. LeCour, Rifleman 82, Rjwilmsi, RobertStar20, Roberta F., Robin Patterson, Robinh, Romanm, Ronhjones, Rursus, SHARD, Saperaud, Saros136, Scarian, SchfiftyThree, Schneelocke, Scot.parker, Scott Adler, Sengkang, Sfuerst, Shaddack, Sheitan, Silentaria, Sir Nicholas de Mimsy-Porpington, Sixpence, Skraz, Sl, Slashme, Smallweed, Snarius, Sound Minded Man, Specter01010, Squids and Chips, SteffenB, Stephenb, Stifynsemons, Stone, Strait, Suisui, Sunborn, Syrthiss, Szaszicska, Tagishsimon, The Hokkaido Crow, The Sanctuary Sparrow, Theeman0000, Theone00, Thricecube, Tide rolls, Tisdalepardi, Tom harrison, Tomcruise0011, Trevor MacInnis, Useight, VASANTH S.N., Vanka5, VasilievVV, Verdatum, Versageek, Versus22, Voyagerfan5761, Vsmith, Vuong Ngan Ha, Watch37264, WereSpielChequers, Weregerbil, West Brom 4ever, Wikiborg, Wilfred Glendon XXVI, Wimt, Wizard191, Wrenchelle, Xxduckyxx, Xy7, Yamaguchi先生, Yettie0711, Yonatan, Yyy, 738 anonymous edits
Image Sources, Licenses and Contributors file:cubic-body-centered.svg Source: http://en.wikipedia.org/w/index.php?title=File:Cubic-body-centered.svg License: GNU Free Documentation License Contributors: User:Stannered file:Electron shell 024 Chromium.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_024_Chromium.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 file:Chrom 1.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Chrom_1.jpg License: Public Domain Contributors: http://de.wikipedia.org/w/index.php?title=Benutzer:Tomihahndorf&action=edit File:Chromit 1.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Chromit_1.jpg License: unknown Contributors: User:Kosioryt Image:Chromium in water pourbiax diagram.png Source: http://en.wikipedia.org/w/index.php?title=File:Chromium_in_water_pourbiax_diagram.png License: Public Domain Contributors: Cadmium, 4 anonymous edits File:Chlorid chromitý.JPG Source: http://en.wikipedia.org/w/index.php?title=File:Chlorid_chromitý.JPG License: Public Domain Contributors: Ondřej Mangl File:Chromium(III)-chloride-purple-anhydrous-sunlight.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Chromium(III)-chloride-purple-anhydrous-sunlight.jpg License: Public Domain Contributors: Ben Mills File:Chrom(VI)-oxid.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Chrom(VI)-oxid.jpg License: GNU Free Documentation License Contributors: BXXXD File:Chroman sodný.JPG Source: http://en.wikipedia.org/w/index.php?title=File:Chroman_sodný.JPG License: Public Domain Contributors: Ondřej Mangl Image:5-fold chromium.png Source: http://en.wikipedia.org/w/index.php?title=File:5-fold_chromium.png License: Public Domain Contributors: User:Rifleman 82 Image:Crocoite from Tasmania.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Crocoite_from_Tasmania.jpg License: unknown Contributors: Eric Hunt File:Cut Ruby.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Cut_Ruby.jpg License: Public Domain Contributors: User:Bkell Image:Chromium - world production trend.svg Source: http://en.wikipedia.org/w/index.php?title=File:Chromium_-_world_production_trend.svg License: Public Domain Contributors: User:Leyo File:World Chromium Production 2002.svg Source: http://en.wikipedia.org/w/index.php?title=File:World_Chromium_Production_2002.svg License: Public Domain Contributors: stone Image:Motorcycle Reflections bw edit.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Motorcycle_Reflections_bw_edit.jpg License: Creative Commons Attribution 2.5 Contributors: User:Atoma File:Laidlaw school bus.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Laidlaw_school_bus.jpg License: Public Domain Contributors: User:Dori
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17
Manganese
1
Manganese chromium ← manganese → iron ↑ Mn ↓ Tc
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
25Mn Periodic table
Appearance silvery metallic
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling
Manganese
2
pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
1228
1347
1493
1691
1955
2333
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1509.0 kJ·mol−1 3rd: 3248 kJ·mol−1Atomic radiusCovalent radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityThermal expansionSpeed of sound (thin rod) Young's modulusBulk modulusMohs hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of manganese iso 52
Mn
N.A. syn
half-life 5.591 d
DM
DE (MeV)
DP
ε
-
52
β+
0.575
52
γ
0.7, 0.9, 1.4
-
Cr Cr
53
syn
3.74 ×106 y
ε
-
53
54
syn
312.3 d
ε
1.377
54
γ
0.834
-
Mn Mn
55
Mn
100%
Cr Cr
55
Mn is stable with 30 neutron
manganese, Mn, 25 transition metal7, 4, d54.938045(5) g·mol−1 [Ar] 4s2 3d5 2, 8, 13, 2 (Image) solid 7.21 g·cm−3 5.95 g·cm−3 1519 K,1246 °C,2275 °F 2334 K,2061 °C,3742 °F 12.91 kJ·mol−1 221 kJ·mol−1 (25 °C) 26.32 J·mol−1·K−17, 6, 5, 4, 3, 2, 1, -1, -2, -3 (oxides: acidic, basic or amphoteric depending on the oxidation state) 1.55 (Pauling scale) 1st: 717.3 kJ·mol−1127 pm139±5 (low spin), 161±8 (high spin) pm cubic paramagnetic (20 °C) 1.44 µΩ·m (300 K) 7.81 W·m−1·K−1 (25 °C) 21.7 µm·m−1·K−1 (20 °C) 5150 m/s 198 GPa 120 GPa 6.0 196 MPa 7439-96-5 Manganese (pronounced /ˈmæŋɡəniːz/) is a chemical element, designated by the symbol Mn. It has the atomic number 25. It is found as a free element in nature (often in combination with iron), and in many minerals. As a free element, manganese is a metal with important industrial metal alloy uses, particularly in stainless steels. Manganese phosphating is used as a treatment for rust and corrosion prevention on steel. Manganese ions have various colors, depending on their oxidation state, and are used industrially as pigments. The permanganates of sodium, potassium and barium are powerful oxidizers. Manganese dioxide is used as the cathode (electron acceptor) material in standard and alkaline disposable dry cells and batteries. Manganese(II) ions function as cofactors for a number of enzymes in higher organisms, where they are essential in detoxification of superoxide free radicals. The element is a required trace mineral for all known living organisms. In larger amounts, and apparently with far greater activity by inhalation, manganese can cause a poisoning syndrome in mammals, with neurological damage which is sometimes irreversible.
Manganese
3
Characteristics Physical Manganese is a gray–white metal, resembling iron. It is a hard metal and is very brittle, fusible with difficulty, but easily oxidized.[1] Manganese metal and its common ions are paramagnetic.[2]
Occurrence See also manganese minerals. Manganese makes up about 1000 ppm (0.1%) of the Earth's crust, making it the 12th most abundant element there.[3] Soil contains 7–9000 ppm of manganese with an average of 440 ppm.[3] Seawater has only 10 ppm manganese and the atmosphere contains 0.01 µg/m3.[3] Manganese occurs principally as pyrolusite (MnO2), braunite, (Mn2+Mn3+6)(SiO12),[4] psilomelane (Ba,H2O)2Mn5O10, and to a lesser extent as Manganese ore rhodochrosite (MnCO3). Land-based resources are large but irregularly distributed. Over 80% of the known world manganese resources are found in South Africa and Ukraine. Other important manganese deposits are in China, Australia, Brazil, Gabon, Ghana, India, and Mexico. In 1978 it was estimated that 500 billion tons of manganese nodules exist on the ocean floor.[5] Attempts to find economically viable methods of harvesting manganese nodules were abandoned in the 1970s.[6]
Isotopes Naturally occurring manganese is composed of 1 stable 55 isotope; Mn. 18 radioisotopes have been 53 54 characterized with the most stable being Mn with a half-life of 3.7 million years, Mn with a half–life of 312.3 days, and 52Mn with a half–life of 5.591 days. All of the remaining radioactive isotopes have half lives that are less than 3 hours and the majority of these have half lives that are less than 1 minute. This element also has 3 meta states.[7] Psilomelane (manganese ore)
Manganese is part of the iron group of elements, which are thought to be synthesized in large stars shortly before the supernova explosion. 53Mn decays to 53Cr with a half-life of 3.7 million years. Because of its relatively short half-life, 53Mn is an extinct radionuclide. Manganese isotopic contents are typically combined with chromium isotopic contents and have found application in isotope geology and radiometric dating. Mn–Cr isotopic ratios reinforce the evidence from 26Al and 107Pd for the early history of the solar system. Variations in 53Cr/52Cr and Mn/Cr ratios from several meteorites indicate an initial 53 Mn/55Mn ratio that suggests Mn–Cr isotopic systematics must result from in–situ decay of 53 Mn in differentiated planetary bodies. Hence 53Mn provides additional evidence for nucleosynthetic processes immediately before coalescence of the solar system.[7]
Manganese
4
The isotopes of manganese range in atomic weight from 46 u (46Mn) to 65 u (65Mn). The primary decay mode before the most abundant stable isotope, 55Mn, is electron capture and the primary mode after is beta decay.[7]
Chemistry and Compounds Oxidation states [8] [9] of manganese 0
Mn2(CO)10
+1
K5[Mn(CN)6NO]
+2
MnCl2
+3
MnF3
+4
MnO2
+5
Na3MnO4
+6
K2MnO4
+7
KMnO4
The most common oxidation states of manganese are +2, +3, +4, +6 and +7, though oxidation states from 0 to +7 are observed. Mn2+ often competes with Mg2+ in biological systems. Manganese compounds where manganese is in oxidation state +7, which are restricted to the oxide Mn2O7 and compounds of the intensely purple permanganate anion MnO4−, are powerful oxidizing agents.[1] . Oxidation states +5 (blue) and +6 (green) are both oxidizing and vulnerable to disproportionation. Methylcyclopentadienyl manganese tricarbonyl is used as an additive in unleaded gasoline to boost octane rating and reduce engine knocking. The manganese in this unusual organometallic compound is in the +1 oxidation state.[10] The most stable oxidation state for manganese is +2, which has a pink to red color, and many manganese(II) compounds are known, such as manganese(II) sulfate The mineral rhodochrosite is (MnSO4) and manganese(II) chloride (MnCl2). This manganese(II) carbonate oxidation state is also seen in the mineral rhodochrosite, (manganese(II) carbonate). The +2 oxidation state is the state used in living organisms for essential functions; all of the other states are much more toxic.
Manganese
5
The +3 oxidation state is known, in compounds such as manganese(III) acetate, but these are quite powerful oxidizing agents and also disproprotionate in solution to Mn(II) and Mn(IV) Solid compounds of Mn(III) are characterized by its preference for distorted octahedral coordination due to the Jahn-Teller effect and its strong purple-red colour.
Manganese(II) chloride
Manganese(IV) oxide (manganese dioxide, MnO2) is used as a reagent in organic chemistry for the oxidation of benzylic alcohols (i.e. adjacent to an aromatic ring). Manganese dioxide has been used since antiquity to oxidatively neutralize the greenish tinge in glass caused by trace amounts of iron contamination.[11] MnO2 is also used in the manufacture of oxygen and chlorine, and in drying black paints. In some preparations it is a brown pigment that can be used to make paint and is a constituent of natural umber. Manganese(IV) oxide was used in the original type of dry cell battery as an electron acceptor from zinc, and is the blackish material found when opening A solution of KMnO4 in water, in a volumetric flask carbon–zinc type flashlight cells. The same material also functions in newer alkaline batteries (usually battery cells), which use the same basic reaction, but a different electrolyte mixture.[12] [13] The oxidation state 5+ can be obtained if manganese dioxide is dissolved in molten sodium nitrite.[14] Manganate (VI) salts can also be produced by dissolving Mn compounds in alkaline melts in air. Permanganate (+7 oxidation state) manganese compounds are purple, and can color glass an amethyst color. Potassium permanganate, sodium permanganate and barium permanganate are all potent oxidizers. Potassium permanganate, also called Condy's crystals, is a commonly used laboratory reagent because of its oxidizing properties and finds use as a topical medicine (for example, in the treatment of fish diseases). Solutions of potassium permanganate were among the first stains and fixatives to be used in the preparation of biological cells and tissues for electron microscopy.[15] Substitutes: Manganese has no satisfactory substitute in its major applications, which are related to metallurgical alloy use.[16] In minor applications, (e.g., manganese phosphating), zinc and sometimes vanadium are viable substitutes. In disposable battery manufacture, standard and alkaline cells using manganese will probably eventually be mostly replaced with lithium battery technology.
Manganese
6
History The origin of the name manganese is complex. In ancient times, two black minerals from Magnesia in what is now modern Greece were both called magnes, but were thought to differ in gender. The male magnes attracted iron, and was the iron ore we now know as lodestone or magnetite, and which probably gave us the term magnet. The female magnes ore did not attract iron, but was used to decolorize glass. This feminine magnes was later called magnesia, known now in modern times as pyrolusite or manganese dioxide. This mineral is never magnetic (although manganese itself is paramagnetic). In the 16th century, the latter compound was called manganesum (note the two n's instead of one) by glassmakers, possibly as a corruption of two words since alchemists and glassmakers eventually had to differentiate a magnesia negra (the black ore) from magnesia alba (a white ore, also from Magnesia, also useful in glassmaking). Michele Mercati called magnesia negra Manganesa, and finally the metal isolated from it became known as manganese (German: Mangan). The name magnesia eventually was then used to refer only to the white magnesia alba (magnesium oxide), which provided the name magnesium for that free element, when it was eventually isolated, much later. [17] Several oxides of manganese, for example manganese dioxide, are abundant in nature and due to color these oxides have been used as since the Stone Age. The cave paintings in Gargas contain manganese as pigments and these cave paintings are 30,000 to 24,000 years old.[19] Manganese compounds where used by Egyptian and Some of the cave painting in Lascaux, France use manganese-based [18] pigments.
Roman glass makers, to either remove color from glass or add color to it.[20] The use as glass makers soap continued through the middle ages until modern times and is evident in 14th century glass from Venice.[11]
Manganese
7
Due to the use in glass making manganese dioxide was available to alchemists the first chemists and was used for experiments. Ignatius Gottfried Kaim (1770) and Johann Glauber ( 17th century) discovered that manganese dioxide can be converted to permanganate, a useful laboratory reagent.[21] By the mid-18th century the Swedish chemist Carl Wilhelm Scheele used manganese dioxide to produce chlorine. First hydrochloric acid, or a mixture of dilute sulfuric acid and sodium chloride was reacted with manganese dioxide, later hydrochloric acid from the Leblanc process was used and the manganese dioxide was recycled by the Weldon process. The production of chlorine and hypochlorite containing bleaching agents was a large consumer of manganese ores. Scheele
and
other
chemists
were
aware
that
manganese dioxide contained a new element, but they were not able to isolate it. Johan Gottlieb Gahn was the first to isolate an impure sample of manganese metal in 1774, by reducing the dioxide with carbon.
Credit for first isolating of manganese is usually given to Johan Gottlieb Gahn
The manganese content of some iron ores used in Greece led to the speculations that the steel produced from that ore contains inadvertent amounts of manganese making the Spartan steel exceptionally hard.[22] Around the beginning of the 19th century, manganese was used in steelmaking and several patents were granted. In 1816, it was noted that adding manganese to iron made it harder, without making it any more brittle. In 1837, British academic James Couper noted an association between heavy exposure to manganese in mines with a form of Parkinson's Disease.[23] In 1912, manganese phosphating electrochemical conversion coatings for protecting firearms against rust and corrosion were patented in the United States, and have seen widespread use ever since.[24] With the invention of the Leclanché cell in 1866 and the subsequent improvement of the batteries containing manganese dioxide as cathodic depolarizer increased the demand of manganese dioxide. Until the introduction of the nickel-cadmium battery and lithium containing batteries most of the batteries on the market contained manganese. The Zinc-carbon battery and the alkaline battery normally use industrially-produced manganese dioxide, because natural occurring manganese dioxide contains impurities. In the 20th century, manganese dioxide has seen wide commercial use as the chief cathodic material for commercial disposable dry cells and dry batteries of both the standard (carbon–zinc) and alkaline type.[25]
Manganese
8
Production The most important manganese ore is pyrolusite (MnO2). Most of the other economically important manganese ores show a close spatial relation to the iron ores.[1] Land-based resources are large but irregularly distributed. Over 80% of the [16] known world manganese Percentage of manganese output in 2006 by countries resources are found in South Africa and Ukraine, other important manganese deposits are in Australia, India, China, Gabon and Brazil.[16] Manganese is mined in South Africa, Australia, China,
Spiegeleisen is an iron alloy with a manganese content of approximately 15 %
Brazil, Gabon, Ukraine, India and Ghana and Kazakhstan.[16] [26] US Import Sources (1998–2001): Manganese ore: Gabon, 70%; South Africa, 10%; Australia, 9%; Mexico, 5%; and other, 6%. Ferromanganese: South Africa, 47%; France, 22%; Mexico, 8%; Australia, 8%; and other, 15%. Manganese contained in all manganese imports: South Africa, 31%; Gabon, 21%; Australia, 13%; Mexico, 8%; and other, 27%.
For the production of ferromanganese the manganese ore are mixed with iron ore and carbon and then reduced either in a blast furnace or in an electric arc furnace.[27] The resulting ferromanganese has a manganese content of 30 to 80 %.[1] Pure manganese used for the production of non-iron alloys is produced by leaching manganese ore with sulfuric acid and a subsequent electrowinning process.[28]
Applications Steel Manganese is essential to iron and steel production by virtue of its sulfur-fixing, deoxidizing, and alloying properties. Steelmaking,[29] including its ironmaking component, has accounted for most manganese demand, presently in the range of 85% to 90% of the total demand.[28] Among a variety of other uses, manganese is a key component of low-cost stainless steel formulations.[26] [30] Small amounts of manganese improve the workability of steel at high temperatures, because it forms a high
British Brodie helmet
Manganese
9
melting sulfide and therefore prevents the formation of a liquid iron sulfide at the grain boundaries. If the manganese content reaches 4% the embrittlement of the steel becomes a dominat feature. The embrittlement decreases at higher manganese concentrations and reaches an acceptable level at 8%. The fact that steel containing 8 to 15% of manganese is [31] [32] cold hardening and can obtain a high tensile strength of up to 863 MPa, steel with 12% manganese was used for the British steel helmets. This steel composition was discovered in 1882 by Robert Hadfield and is still known as Hadfield steel.[33]
Aluminium alloys The second large application for manganese is as alloying agent for aluminium. Aluminium with a manganese content of roughly 1.5 % has an increased resistance against corrosion due to the formation grains absorbing impurities which would lead to galvanic corrosion.[34] The corrosion resistant aluminium alloy 3004 and 3104 with a manganese content of 0.8 to 1.5 % are the alloy used for most of the beverage cans.[35] For years prior to 2000 in excess of 1.6 million metric tons have been used of those alloys, with a content of 1 % of manganese this amount would need 16000 metric tons of manganese.[35]
Other use A large amount of manganese dioxide is produced for the use as depolarizer in Zinc-carbon battery and the alkaline battery.[25] In 2002 more than 230,000 tons of manganese dioxide were used for this purpose.[28] The manganese dioxide is reduced to the manganese oxide-hydroxide MnO(OH) during discharging, preventing the formation of hydrogen at the anode of the battery.[12] MnO + H O + e- → MnO(OH) + OH2
2
The metal is very occasionally used in coins; the only United Wartime nickel made from a States coins to use manganese were the "wartime" nickel [36] copper silver manganese alloy from 1942–1945. Due to shortage of raw materials in the war the nickel in the alloy (75 % copper and 25 % nickel) used for the production of the nickel before was substituted by the less critical metals silver and manganese (56 % copper, 35 % silver and 9 % manganese). Since 2000 dollar coins, for example the Sacagawea dollar and the Presidential $1 Coins, are made from a brass containing 7% of manganese with a pure copper core.[37]
Manganese compounds have been used as pigments and for the coloring of ceramics and glass for a long time and still the brown color of ceramic is sometimes based on manganese compounds.[38] In the glass industry two effects of manganese compounds are used. Manganese(III) reacts with iron(II). The reaction induces a strong green color in glass by forming less-colored iron(III) and slightly pink manganese(II), compensating the residual color of the iron(III).[11] Larger amounts of manganese are used to produce pink colored glass.
Manganese
Biological role Manganese is an essential trace nutrient in all forms of life.[39] The classes of enzymes that have manganese cofactors are very broad and include such classes as oxidoreductases, transferases, hydrolases, lyases, isomerases, ligases, lectins, and integrins. The reverse transcriptases of many retroviruses (though not lentiviruses such as HIV) Reactive center of arginase with boronic acid inhibitor. The contain manganese. The best managanese is shown in yellow known manganese-containing polypeptides may be arginase, the diphtheria toxin, and Mn-containing superoxide dismutase (Mn-SOD).[40] Mn-SOD is the type of SOD present in eukaryotic mitochondria, and also in most bacteria (this fact is in keeping with the bacterial-origin theory of mitochondria). The Mn-SOD enzyme is probably one of the most ancient, for nearly all organisms living in the presence of oxygen use it to deal with the toxic effects of superoxide, formed from the 1-electron reduction of dioxygen. Exceptions include a few kinds of bacteria such as Lactobacillus plantarum and related lactobacilli, which use a different non-enzymatic mechanism, involving manganese (Mn2+) ions complexed with polyphosphate directly for this task, indicating how this function possibly evolved in aerobic life. The human body contains about 10 mg of manganese, which is stored mainly in the liver and kidneys. In the human brain the manganese is bound to manganese metalloproteins most notable glutamine synthetase in astrocytes.[41] Manganese is also important in photosynthetic oxygen evolution in chloroplasts in plants. The oxygen evolving complex (OEC) is a part of Photosystem II contained in the thylakoid membranes of chloroplasts; it is responsible for the terminal photooxidation of water during the light reactions of photosynthesis and has a metalloenzyme core containing four atoms of manganese.[42] For this reason, most broad-spectrum plant fertilizers contain manganese.
Precautions Manganese compounds are less toxic than those of other widespread metals such as nickel and copper.[43] However, exposure to manganese dusts and fumes should not exceed the ceiling value of 5 mg/m3[44] even for short periods because of its toxicity level. Manganese poses a particular risk for children due to its propensity to bind to CH-7 receptors. Manganese poisoning has been linked to impaired motor skills and cognitive disorders.[45] The permanganate exhibits a higher toxicity than the manganese(II) compounds. Several fatal intoxications have occurred, although the fatal dose is around 10 g. The strong oxidative effect leads to necrosis of the mucous membrane. For example, the esophagus is affected if the permanganate is swallowed. Only a limited amount is absorbed by the
10
Manganese intestines but this small amount shows the severe effects on the kidneys and on the liver.[46] [47] In 2005, a study suggested a possible link between manganese inhalation and central nervous system toxicity in rats.[48] It is hypothesized that long-term exposure to the naturally occurring manganese in shower water puts up to 8.7 million Americans at risk.[49] [50] [51]
A form of neurodegeneration[52] similar to Parkinson's Disease called "manganism" has been linked to manganese exposure amongst miners and smelters since the early 19th Century.[53] Allegations of inhalation-induced manganism have been made regarding the welding industry. Manganese exposure in United States is regulated by Occupational Safety and Health Administration. [54]
See also • Parkerize
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11
Manganese [16] Corathers, Lisa A. (2009). " Mineral Commodity Summaries 2009: Manganese (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ manganese/ mcs-2009-manga. pdf)" (PDF). United States Geological Survey. . Retrieved 2009-04-30. [17] " Chromium and Manganese (http:/ / www. du. edu/ ~jcalvert/ phys/ chromang. htm)". 2003-01-24. . Retrieved 2009-04-30. [18] Chalmin, Emilie; Menu, Michel; Vignaud, Colette (2003). "Analysis of rock art painting and technology of Palaeolithic painters". Measurement Science and Technology 14: 1590–1597. doi: 10.1088/0957-0233/14/9/310 (http:/ / dx. doi. org/ 10. 1088/ 0957-0233/ 14/ 9/ 310). [19] Chalmin, E.; Vignaud, C.; Salomon, H.; Farges, F.; Susini, J.; Menu, M. (2006). "Minerals discovered in paleolithic black pigments by transmission electron microscopy and micro-X-ray absorption near-edge structure". Applied Physics A 83: 213–218. doi: 10.1007/s00339-006-3510-7 (http:/ / dx. doi. org/ 10. 1007/ s00339-006-3510-7). [20] Sayre, E. V.; Smith, R. W. (Jun 1961). "Compositional Categories of Ancient Glass.". Science 133 (3467): 1824–1826. doi: 10.1126/science.133.3467.1824 (http:/ / dx. doi. org/ 10. 1126/ science. 133. 3467. 1824). PMID 17818999. [21] Rancke-Madsen, E. (1975). "The Discovery of an Element". Centaurus 19 (4): 299–313. doi: 10.1111/j.1600-0498.1975.tb00329.x (http:/ / dx. doi. org/ 10. 1111/ j. 1600-0498. 1975. tb00329. x). [22] Alessio, L; Campagna, M; Lucchini, R (Nov 2007). "From lead to manganese through mercury: mythology, science, and lessons for prevention.". American journal of industrial medicine 50 (11): 779–787. doi: 10.1002/ajim.20524 (http:/ / dx. doi. org/ 10. 1002/ ajim. 20524). [23] Couper, J. (1837). "On the effects of black oxide of manganese when inhaled into the lungs". Br. Ann. Med. Pharmacol. 1: 41–42. [24] Olsen, Sverre E.; Tangstad, Merete; Lindstad, Tor (2007). "History of manganese". Production of Manganese Ferroalloys. Tapir Academic Press. pp. 11–12. ISBN 9788251921916. [25] Preisler, Eberhard (1980). "Moderne Verfahren der Großchemie: Braunstein". Chemie in unserer Zeit 14: 137–148. doi: 10.1002/ciuz.19800140502 (http:/ / dx. doi. org/ 10. 1002/ ciuz. 19800140502). [26] Corathers, Lisa A. (June 2008). " 2006 Minerals Yearbook: Manganese (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ manganese/ myb1-2006-manga. pdf)" (PDF). Washington, D.C.: United States Geological Survey. . Retrieved 2009-04-30. [27] Corathers, L. A.; Machamer, J. F. (2006). " Manganese (http:/ / books. google. de/ books?id=zNicdkuulE4C& pg=PA631)". Industrial Minerals & Rocks: Commodities, Markets, and Uses (7th ed.). SME. pp. 631–636. ISBN 9780873352338. . [28] Zhang, Wensheng; Cheng, Chu Yong (2007). "Manganese metallurgy review. Part I: Leaching of ores/secondary materials and recovery of electrolytic/chemical manganese dioxide". Hydrometallurgy 89: 137–159. doi: 10.1016/j.hydromet.2007.08.010 (http:/ / dx. doi. org/ 10. 1016/ j. hydromet. 2007. 08. 010). [29] Verhoeven., John D. (2007). Steel metallurgy for the non-metallurgist. Materials Park, Ohio: ASM International. pp. 56–57. ISBN 9780871708588. [30] Dastur, Y. N. (1981). "Mechanism of work hardening in Hadfield manganese steel". Metallurgical Transactions A 12: 749. doi: 10.1007/BF02648339 (http:/ / dx. doi. org/ 10. 1007/ BF02648339). [31] Stansbie, John Henry (2007). Iron and Steel (http:/ / books. google. com/ books?id=FyogLqUxW1cC& pg=PA351). Read Books. pp. 351–352. ISBN 9781408626160. . [32] Brady, George S.; Clauser; Henry R.; Vaccari. John A. (2002). Materials handbook : an encyclopedia for managers, technical professionals, purchasing and production managers, technicians, and supervisors (http:/ / books. google. com/ books?id=vIhvSQLhhMEC& pg=PA585). New York, NY: McGraw-Hill. pp. 585–587. ISBN 9780071360760. . [33] Tweedale, Geoffrey (1985). " Sir Robert Abbott Hadfield F.R.S. (1858-1940), and the Discovery of Manganese Steel Geoffrey Tweedale (http:/ / www. jstor. org/ stable/ 531536)". Notes and Records of the Royal Society of London 40 (1): 63–74. doi: 10.1098/rsnr.1985.0004 (http:/ / dx. doi. org/ 10. 1098/ rsnr. 1985. 0004). . [34] " chemical properties of 2024 aluminum allow (http:/ / www. suppliersonline. com/ propertypages/ 2024. asp)". Metal Suppliers Online, LLC.. . Retrieved 2009-04-30. [35] Kaufman, John Gilbert (2000). " Applications for Aluminium Alloys and Tempers (http:/ / books. google. de/ books?id=idmZIDcwCykC& pg=PA93)". Introduction to aluminum alloys and tempers. ASM International. pp. 93–94. ISBN 9780871706898. . [36] Kuwahara, Raymond T.; Skinner III, Robert B.; Skinner Jr., Robert B. (2001). " Nickel coinage in the United States (http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?artid=1071501)". Western Journal of Medicine 175 (2): 112–114. doi: 10.1136/ewjm.175.2.112 (http:/ / dx. doi. org/ 10. 1136/ ewjm. 175. 2. 112). PMID 11483555. . [37] Design of the Sacagawea dollar (http:/ / www. usmint. gov/ mint_programs/ golden_dollar_coin/ index. cfm?action=sacDesign). United States Mint. . Retrieved 2009-05-04.
12
Manganese [38] Shepard, Anna Osler (1956). "Manganese and Iron–Manganese Paints". Ceramics for the archaeologist. Carnegie Institution of Washington. pp. 40–42. ISBN 9780872796201. [39] Emsley, John (2001). "Manganese". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, UK: Oxford University Press. pp. 249–253. ISBN 0198503407. [40] Law, N. (1998). Manganese Redox Enzymes and Model Systems: Properties, Structures, and Reactivity. 46. pp. 305. doi: 10.1016/S0898-8838(08)60152-X (http:/ / dx. doi. org/ 10. 1016/ S0898-8838(08)60152-X). [41] Takeda, A. (2003). "Manganese action in brain function". Brain Research Reviews 41: 79. doi: 10.1016/S0165-0173(02)00234-5 (http:/ / dx. doi. org/ 10. 1016/ S0165-0173(02)00234-5). [42] Dismukes, G. Charles; Willigen, Rogier T. van (2006). "Manganese: The Oxygen-Evolving Complex & Models". Encyclopedia of Inorganic Chemistry. doi: 10.1002/0470862106.ia128 (http:/ / dx. doi. org/ 10. 1002/ 0470862106. ia128). [43] Hasan, Heather (2008). Manganese (http:/ / books. google. de/ books?id=nRmpEaudmTYC& pg=PA31). The Rosen Publishing Group. pp. 31. ISBN 9781404214088. . [44] " Manganese Chemical Background (http:/ / www. environmentwriter. org/ resources/ backissues/ chemicals/ manganese. htm)". Metcalf Institute for Marine and Environmental Reporting University of Rhode Island. 2006-04. . Retrieved 2008-04-30. [45] " Risk Assessment Information System Toxicity Summary for Manganese (http:/ / rais. ornl. gov/ tox/ profiles/ mn. shtml)". Oak Ridge National Laboratory. . [46] Ong, K. L. (1997). "Potassium permanganate poisoning--a rare cause of fatal self poisoning.". Emergency Medicine Journal 14: 43. doi: 10.1136/emj.14.1.43 (http:/ / dx. doi. org/ 10. 1136/ emj. 14. 1. 43). PMID 9023625. [47] Young, R. (1996). "Fatal acute hepatorenal failure following potassium permanganate ingestion". Human & Experimental Toxicology 15: 259. doi: 10.1177/096032719601500313 (http:/ / dx. doi. org/ 10. 1177/ 096032719601500313). PMID 8839216. [48] Elsner, Robert J. F.; Spangler, John G. (2005). "Neurotoxicity of inhaled manganese: Public health danger in the shower?". Medical Hypotheses 65 (3): 607–616. doi: 10.1016/j.mehy.2005.01.043 (http:/ / dx. doi. org/ 10. 1016/ j. mehy. 2005. 01. 043). [49] Elsner, Rj; Spangler, Jg (2005). "Neurotoxicity of inhaled manganese: public health danger in the shower?". Medical hypotheses 65 (3): 607–616. doi: 10.1016/j.mehy.2005.01.043 (http:/ / dx. doi. org/ 10. 1016/ j. mehy. 2005. 01. 043). ISSN 0306-9877 (http:/ / worldcat. org/ issn/ 0306-9877). PMID 15913899. [50] Finley, John Weldon (1999). "Manganese deficiency and toxicity: Are high or low dietary amounts of manganese cause for concern?". BioFactors 10: 15. doi: 10.1002/biof.5520100102 (http:/ / dx. doi. org/ 10. 1002/ biof. 5520100102). [51] Barceloux, Donald (1999). "Manganese". Clinical Toxicology 37: 293. doi: 10.1081/CLT-100102427 (http:/ / dx. doi. org/ 10. 1081/ CLT-100102427). [52] Normandin, Louise (2002). Metabolic Brain Disease 17: 375. doi: 10.1023/A:1021970120965 (http:/ / dx. doi. org/ 10. 1023/ A:1021970120965). [53] Crossgrove, J; Zheng, W (Dec 2004). "Manganese toxicity upon overexposure.". NMR in biomedicine 17 (8): 544–553. doi: 10.1002/nbm.931 (http:/ / dx. doi. org/ 10. 1002/ nbm. 931). ISSN 0952-3480 (http:/ / worldcat. org/ issn/ 0952-3480). PMID 15617053. [54] " Safety and Health Topics: Manganese Compounds (as Mn) (http:/ / www. osha. gov/ dts/ chemicalsampling/ data/ CH_250190. html)". . 080225 osha.gov
External links • National Pollutant Inventory – Manganese and compounds Fact Sheet (http:/ / www. npi. gov. au/ database/ substance-info/ profiles/ 52. html) • WebElements.com – Manganese (http:/ / www. webelements. com/ webelements/ elements/ text/ Mn/ index. html) • International Manganese Institute (http:/ / www. manganese. org) • Neurotoxicity of inhaled manganese: Public health danger in the shower? (http:/ / dx. doi. org/ 10. 1016/ j. mehy. 2005. 01. 043) • NIOSH Manganese Topic Page (http:/ / www. cdc. gov/ niosh/ topics/ manganese/ ) • Link Found Between Parkinson's Disease Genes And Manganese Poisoning (http:/ / www. sciencedaily. com/ releases/ 2009/ 02/ 090201141559. htm)
13
Article Sources and Contributors
Article Sources and Contributors Manganese Source: http://en.wikipedia.org/w/index.php?oldid=307659238 Contributors: 2D, Abihail, Adamrush, Adrian.benko, Ahoerstemeier, Akusarujin, Alansohn, Anabus, Andreala, Andres, Ann Stouter, Anonymous Dissident, Archimerged, ArglebargleIV, Atjesse, Aussie Alchemist, Baby Jane, BananaFiend, Bassbonerocks, Bcorr, Beetstra, Benbest, Bender235, Bendzh, Benjah-bmm27, Betacommand, Biochemnick, BlueEarth, Bobo192, Bomac, Bork, Brian0918, Brinerustle, Bryan Derksen, Btyner, Bubb13z, Cadmium, Caltas, Capricorn42, Carnildo, Catskul, Chameleon, Charles Matthews, Cheese4, Chill doubt, Chris 73, Clarityfiend, Coemgenus, Colo1115, Conversion script, Coppertwig, Csladic, Cssiitcic, Dakirw8, Daniel bg, Darrien, David Latapie, Davidjk, Dcljr, Debresser, Delirium, Delta G, Dennis Brown, DerHexer, Dogsgomoo, Donarreiskoffer, Dwmyers, EDM, Edgar181, Eeekster, El C, Eldin raigmore, Electron9, Emiley Bee, Emperorbma, Eog1916, Epbr123, Evand, Farosdaughter, Femto, Finlay McWalter, Fredrik, Frosty0814snowman, Fusionmix, GSlicer, Garden, Gbr3, Gekedo, Gman124, Gogo Dodo, Graingert, Grendelkhan, Hak-kâ-ngìn, Hephaestos, Iapetus, Icairns, Ice Czar, Ideyal, Ixfd64, J.delanoy, J14fusion, JCarriker, Jalscott, Jamesontai, Jaraalbe, Jaronharding02, Jerzy, Jjdon, Joanjoc, John, John254, JohnKoziar, Jons63, Jose77, Jpeloquin, Jpotherington, Juliancolton, Julianp, Junglecat, KF, KJS77, Kaiba, Karl-Henner, Karlhahn, Keilana, Knaggs, Kormoran, Kthb, Kwamikagami, LA2, LanceBarber, LarryMorseDCOhio, LeaveSleaves, Leuko, Luk, Lumos3, MZMcBride, Mani1, Marek69, Markjoseph125, Markkawika, Materialscientist, Mav, Mdf, Megaman en m, Mgimpel, Michael Hardy, MindstormsKid, Minesweeper, Mintrick, Mkweise, Modernist, Mr Rookles, NJPharris, Naddy, Nahum Reduta, Nakon, Nergaal, NewEnglandYankee, Nick Y., Nihiltres, Nosivad, Osip7315, Oxymoron83, Paraballo, Pb30, Philip Trueman, Plantsurfer, Poolkris, Pras, Prolog, Psyche825, RTC, Ramius, Ranveig, Rawling, Raymondwinn, RexNL, Roberta F., Romanm, Sally M. Peters, Saperaud, Sbharris, Sceptre, Schneelocke, Sengkang, Shablog, Shaddack, Shafei, Shappy, Sheitan, Sl, Soher23, Squids and Chips, Srnec, Statsone, Stephenb, Stifynsemons, Stone, Suisui, Sunborn, T Tom The Bomb, THEN WHO WAS PHONE?, Tagishsimon, The Obento Musubi, The Rambling Man, TheAdventMaster, TheSeer, Thortveitite, Thricecube, Tisdalepardi, Tivedshambo, Tornvmax, Trevor MacInnis, Uirauna, Until It Sleeps, Uppland, Veinor, Vsmith, Vuo, Walkerma, WarthogDemon, Watch37264, Werdan7, WhatamIdoing, Whosasking, Wiki Raja, Wilson44691, Wimt, Wmahan, Yaf, Yamakiri, Yamamoto Ichiro, Yyy, 501 anonymous edits
Image Sources, Licenses and Contributors file:cubic.svg Source: http://en.wikipedia.org/w/index.php?title=File:Cubic.svg License: GNU Free Documentation License Contributors: User:Stannered file:Electron shell 025 Manganese.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_025_Manganese.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 file:Mangan 1-crop.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Mangan_1-crop.jpg License: GNU Free Documentation License Contributors: Original uploader was Tomihahndorf at de.wikipedia Image:ManganeseOreUSGOV.jpg Source: http://en.wikipedia.org/w/index.php?title=File:ManganeseOreUSGOV.jpg License: unknown Contributors: Duesentrieb, Joanjoc, Saperaud, רדמ לבוי Image:Mineraly.sk - psilomelan.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Mineraly.sk_-_psilomelan.jpg License: unknown Contributors: Helix84, Saperaud File:The Searchlight Rhodochrosite Crystal.jpg Source: http://en.wikipedia.org/w/index.php?title=File:The_Searchlight_Rhodochrosite_Crystal.jpg License: unknown Contributors: Eric Hunt File:Chlorid manganatý.JPG Source: http://en.wikipedia.org/w/index.php?title=File:Chlorid_manganatý.JPG License: Public Domain Contributors: Ondřej Mangl Image:KMnO4 in H2O.jpg Source: http://en.wikipedia.org/w/index.php?title=File:KMnO4_in_H2O.jpg License: unknown Contributors: David Mülheims (David Mülheims, Germany) File:Lascaux-aurochs.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Lascaux-aurochs.jpg License: unknown Contributors: Some caveman/cavemen. File:Gahn Johan Gottlieb.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Gahn_Johan_Gottlieb.jpg License: Public Domain Contributors: Weeks, Mary Elvira File:World Manganese Production 2006.svg Source: http://en.wikipedia.org/w/index.php?title=File:World_Manganese_Production_2006.svg License: Public Domain Contributors: stone File:Spiegeleisen.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Spiegeleisen.jpg License: unknown Contributors: Tillman File:M1917helmet.jpg Source: http://en.wikipedia.org/w/index.php?title=File:M1917helmet.jpg License: GNU Free Documentation License Contributors: Joshua R. Murray File:War Nickle.jpg Source: http://en.wikipedia.org/w/index.php?title=File:War_Nickle.jpg License: unknown Contributors: Cholmes75, Stauba File:Arginase.jpeg Source: http://en.wikipedia.org/w/index.php?title=File:Arginase.jpeg License: unknown Contributors: Eswiss, Materialscientist
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14
Iron
1
Iron manganese ← iron → cobalt ↑ Fe ↓ Ru
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HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
26Fe Periodic table
Appearance lustrous metallic with a grayish tinge
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
1728
1890
2091
2346
2679
3132
Atomic properties ElectronegativityIonization energies (more) 2nd: 1561.9 kJ·mol−1 3rd: 2957 kJ·mol−1Atomic radiusCovalent radius Miscellaneous Magnetic ordering 1043 K Electrical resistivityThermal conductivityThermal
Iron
2
expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusPoisson ratioMohs hardnessVickers hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of iron iso
N.A.
half-life
DM
DE (MeV)
DP
54
5.8%
>3.1×1022y
2ε capture
?
54
55
syn
2.73 y
ε capture
0.231
55
56
91.72%
56
57
2.2%
57
58
0.28%
58
59
syn
44.503 d
β−
1.565
59
60
syn
1.5×106 y
β−
3.978
60
Fe Fe Fe Fe Fe Fe Fe
Cr Mn
Fe is stable with 30 neutron Fe is stable with 31 neutron Fe is stable with 32 neutron Co Co
iron, Fe, 26 transition metal8, 4, d55.845(2) g·mol−1 [Ar] 3d6 4s2 2, 8, 14, 2 (Image) solid 7.874 g·cm−3 6.98 g·cm−3 1811 K,1538 °C,2800 °F 3134 K,2862 °C,5182 °F 13.81 kJ·mol−1 340 kJ·mol−1 (25 °C) 25.10 J·mol−1·K−1 1.83 (Pauling scale) 1st: 762.5 kJ·mol−1 126 pm 132±3 (low spin), 152±6 (high spin) pm ferromagnetic (20 °C) 96.1 nΩ·m (300 K) 80.4 W·m−1·K−1 (25 °C) 11.8 µm·m−1·K−1 (r.t.) (electrolytic) 5120 m·s−1 211 GPa 82 GPa 170 GPa 0.29 4.0 608 MPa 490 MPa 7439-89-6 Iron (pronounced /ˈаɪ.ərn/) is a chemical element with the symbol Fe (Latin: ferrum) and atomic number 26. Iron is a group 8 and period 4 element. Iron and iron alloys (steels) are by far the most common metals and the most common ferromagnetic materials in everyday use. Fresh iron surfaces are lustrous and silvery-grey in colour, but oxidise in air to form a red or brown coating of ferrous oxide or rust. Pure single crystals of iron are soft (softer than aluminium), and the addition of minute amounts of impurities, such as carbon, significantly strengthens them. Alloying iron with appropriate small amounts (up to a few per cent) of other metals and carbon produces steel, which can be 1,000 times harder than pure iron. Iron-56 is the heaviest stable isotope produced by the alpha process in stellar nucleosynthesis; heavier elements than iron and nickel require a supernova for their formation. Iron is the most abundant element in the core of red giants, and is the most abundant metal in iron meteorites and in the dense metal cores of planets such as Earth.
Characteristics Pure iron is a metal but is rarely found in this form on the surface of the earth because it oxidizes readily in the presence of oxygen and moisture. In order to obtain metallic iron, oxygen must be removed from naturally occurring ores by chemical reduction – mainly of the iron ore hematite (Fe2O3) by carbon at high temperature. The properties of iron can be modified by alloying it with various other metals (and some non-metals, notably carbon and silicon) to form steels. Nuclei of iron atoms have some of the highest binding energies per nucleon, surpassed only by the nickel isotope 62Ni. The universally most abundant of the highly stable nuclides is, however, 56Fe. This is formed by nuclear fusion in stars. Although a further tiny energy gain could be extracted by synthesizing 62Ni, conditions in stars are unsuitable for this
Iron
3
process to be favoured. Elemental distribution on Earth greatly favours iron over nickel, and also presumably in supernova element production.[1] Iron (as Fe2+, ferrous ion) is a necessary trace element used by almost all living organisms. The only exceptions are several organisms that live in iron-poor environments and have evolved to use different elements in their metabolic processes, such as manganese instead of iron for catalysis, or hemocyanin instead of hemoglobin. Iron-containing enzymes, usually containing heme prosthetic groups, participate in catalysis of oxidation reactions in biology, and in transport of a number of soluble gases. See hemoglobin, cytochrome, and catalase.
Mechanical properties Characteristic values of tensile strength (TS) and Brinell hardness (BH) of different forms of iron. Material
TS (MPa)
BH (Brinell)
Iron whiskers
11000
Ausformed (hardened) steel
2930
850-1200
Martensitic steel
2070
600
Bainitic steel
1380
400
Pearlitic steel
1200
350
Cold-worked iron
690
200
Small-grain iron
340
100
Iron containing dissolved carbon
140
40
Single crystal of pure iron
10
3
Mechanical properties of iron and its alloys are traditionally evaluated using various measurements, such as Brinell test, Rockwell test, tensile strength and other; their results are so much consistent among each other that universal relations are often used to relate results of one measurement to another.[] [2] Those measurements reveal that mechanical properties of iron crucially depend on purity: Purest research-purpose single crystals of iron are softer than aluminium. Addition of only 10 parts per million of carbon doubles their strength.[] The hardness increases rapidly with carbon content up to 0.2% and saturates at ~0.6%.[3] The purest industrially produced iron (about 99.99% purity) has hardness of 20-30 Brinell [4]
Allotropes Iron represents perhaps the best-known example of allotropy in a metal. There are three allotropic forms of iron, known as α, γ and δ. As molten iron cools down it crystallizes at 1538 °C into its δ allotrope, which has a body-centred cubic (bcc) crystal structure. As it cools further its crystal structure changes to face-centred cubic (fcc) at 1394 °C, when it is known as γ-iron, or austenite. At 912 °C the crystal structure again becomes bcc as α-iron, or ferrite, is formed, and at 770 °C (the Curie point, Tc) the iron becomes magnetic. As the iron passes through the Curie temperature there is no change in crystalline structure, but there is a change in "domain structure", where each domain contains iron atoms with a particular electronic spin. In
Iron
4
unmagnetized iron, all the electronic spins of the atoms within one domain are in the same direction; however, in neighbouring domains they point in various directions and thus cancel out. In magnetized iron, the electronic spins of all the domains are all aligned, so that the magnetic effects of neighbouring domains reinforce each other. Although each domain contains billions of atoms, they are very small, about 10 microns across. Iron is of most importance when mixed with certain other metals and with carbon to form steels. There are many types of steels, all with different properties; and an understanding of the properties of the allotropes of iron is key to the manufacture of good quality steels. Alpha iron, also known as ferrite, is the most stable form of iron at normal temperatures. It is a fairly soft metal that can dissolve only a small concentration of carbon (no more than 0.021% by mass at 910 °C).[5] Above 912 °C and up to 1400 °C α-iron undergoes a phase transition from body-centred cubic to the face-centred cubic configuration of γ-iron, also called austenite. This is similarly soft and metallic but can dissolve considerably more carbon (as much as 2.04% by mass at 1146 °C). This form of iron is used in the type of stainless steel used for making cutlery, and hospital and food-service equipment.
Occurrence Iron is the sixth most abundant element in the Universe, formed as the final act of nucleosynthesis, by silicon fusing in massive stars. While it makes up about 5% of the Earth's crust, the Earth's core is believed to consist largely of an iron-nickel alloy constituting 35% of the mass of the Earth as a whole. Iron is consequently the most abundant element on Earth, but only the fourth most abundant element in the Earth's crust.[6] Most of the iron in the crust is found combined with oxygen as iron oxide minerals such as hematite and magnetite.
The red appearance of this water is due to ferric ion, Iron(III) or Fe3+, in the rocks.
About 1 in 20 meteorites consist of the unique iron-nickel minerals taenite (35–80% iron) and kamacite (90–95% iron). Although rare, iron meteorites are the major form of natural metallic iron on the Earth's surface. The red colour of the surface of Mars is thought to derive from an iron oxide-rich regolith.
Iron
Isotopes Naturally occurring iron consists of four isotopes: 5.845% of radioactive 54Fe (half-life: >3.1×1022 years), 91.754% of stable 56Fe, 2.119% of stable 57Fe and 0.282% of stable 58 Fe. 60Fe is an extinct radionuclide of long half-life (1.5 million years). Much of the past work on measuring the isotopic composition of Fe has centred on determining 60Fe variations due to processes accompanying nucleosynthesis (i.e., meteorite studies) and ore formation. In the last decade however, advances in mass spectrometry technology have allowed the detection and quantification of minute, naturally occurring variations in the ratios of the stable isotopes of iron. Much of this work has been driven by the Earth and planetary science communities, although applications to biological and industrial systems are beginning to emerge.[7] The most abundant iron isotope 56Fe is of particular interest to nuclear scientists. A common misconception is that this isotope represents the most stable nucleus possible, and that it thus would be impossible to perform fission or fusion on 56Fe and still liberate energy. This is not true, as both 62Ni and 58Fe are more stable, being the most stable nuclei. However, since 56Ni is much more easily produced from lighter nuclei in the alpha process in nuclear reactions in supernovae (see silicon burning process), nickel-56 (14 alpha particles) is the endpoint of fusion chains inside extremely massive stars, since addition of another alpha would result in zinc-60, which requires a great deal more energy. This nickel-56, which has a half-life of about 6 days, is therefore made in quantity in these stars, but soon decays by two successive positron emissions within supernova decay products in the supernova remnant gas cloud, to first radioactive cobalt-56, and then stable iron-56. This last nuclide is therefore common in the universe, relative to other stable metals of approximately the same atomic weight. In phases of the meteorites Semarkona and Chervony Kut a correlation between the concentration of 60Ni, the daughter product of 60Fe, and the abundance of the stable iron isotopes could be found which is evidence for the existence of 60Fe at the time of formation of the solar system. Possibly the energy released by the decay of 60Fe contributed, together with the energy released by decay of the radionuclide 26Al, to the remelting and differentiation of asteroids after their formation 4.6 billion years ago. The abundance of 60 Ni present in extraterrestrial material may also provide further insight into the origin of the solar system and its early history. Of the stable isotopes, only 57Fe has a nuclear spin (−1/2).
Chemistry and compounds Iron forms compounds mainly in the +2 and +3 oxidation states. Traditionally, iron(II) compounds have been called ferrous, and iron(III) compounds ferric. There are many compounds in each of the oxidation states (see Category:Iron compounds for a list); representative examples would include iron(II) sulfate (FeSO4) and iron(III) chloride (FeCl3). There are also numerous examples of compounds that contain iron atoms in both of these oxidation states, such as magnetite and prussian blue. The ferrate anion [FeO4]2contains an iron(VI) centre, its highest known oxidation state, and is present, for example in potassium ferrate (K2FeO4). There are numerous organometallic compounds (such as iron pentacarbonyl) that contain formally zerovalent (or lower) iron.
5
Iron
6
History The first wrought iron used by mankind during prehistory came from meteors. The smelting of iron in bloomeries began in the second millennium BC. Artefacts from smelted iron occur in India from 1800-1200 BC.[8] in the Levant from about 1500 BC (suggesting smelting in Anatolia or the Caucasus).[9] The symbol for Mars has been used since ancient times to represent iron.
[10]
Cast iron was first produced in China about 550 BC,[11] but not in Europe until the medieval period. During the medieval period, means were found in Europe of producing wrought iron from cast iron (in this context known as pig iron) using finery forges. For all these processes, charcoal was required as fuel. Steel (with smaller carbon content than pig iron but more than wrought iron) was first produced in antiquity. New methods of producing it by carburizing bars of iron in the cementation process were devised in the 17th century AD. In the Industrial Revolution, new methods of producing bar iron without charcoal were devised and these were later applied to produce steel. In the late 1850s, Henry Bessemer invented a new steelmaking process, involving blowing air through molten pig iron, to produce mild steel. This and other 19th century and later processes have led to wrought iron no longer being produced.
The Delhi iron pillar is an example of the iron extraction and processing methodologies of India. The iron pillar at Delhi has withstood corrosion for the last 1600 years.
Iron
7
Industrial production The production of iron or steel is a process unless the desired final product is cast iron. The first stage is to produce pig iron in a blast furnace. The second is to make wrought iron or steel from pig iron by a further process.
Blast furnace Ninety percent of all mining of metallic ores is for the extraction of iron. Industrially, iron is produced starting Iron ore pellets from Kiruna, Sweden. from iron ores, principally hematite (nominally Fe2O3) and magnetite (Fe3O4) by a carbothermic reaction (reduction with carbon) in a blast furnace at temperatures of about 2000 °C. In a blast furnace, iron ore, carbon in the form of coke, and a flux such as limestone (which is used to remove impurities in the ore which would otherwise clog the furnace with solid material) are fed into the top of the furnace, while a blast of heated air is forced into the furnace at the bottom. In the furnace, the coke reacts with oxygen in the air blast to produce carbon monoxide: 2 C + O → 2 CO 2
The carbon monoxide reduces the iron ore (in the Iron output in 2005
chemical equation below, hematite) to molten iron, becoming carbon dioxide in the process:
3 CO + Fe2O3 → 2 Fe + 3 CO2 The flux is present to melt impurities in the ore, principally silicon dioxide sand and other silicates. Common fluxes include limestone (principally calcium carbonate) and dolomite (calcium-magnesium carbonate). Other fluxes may be used depending on the impurities that need to be removed from the ore. In the heat of the furnace the limestone flux decomposes to calcium oxide (quicklime): CaCO3 → CaO + CO2 Then calcium oxide combines with silicon dioxide to form a slag. CaO + SiO2 → CaSiO3 The slag melts in the heat of the furnace. In the bottom of the furnace, the molten slag floats on top of the denser molten iron, and apertures in the side of the furnace are opened to run off the iron and the slag separately. The iron once cooled, is called pig iron, while the slag can be used as a material in road construction or to improve mineral-poor soils for agriculture.
Iron
8
In 2005, approximately 1,544 Mt (million metric tons) of iron ore were produced worldwide. China was the top producer of iron ore with at least one-fourth world share followed by Brazil, Australia and India, reports the British Geological Survey.
Further processes Pig iron is not pure iron, but has 4-5% carbon dissolved in it with small amounts of other impurities like sulfur, magnesium, phosphorus and manganese. As the carbon is the major impurity, the iron (pig iron) becomes brittle and hard. This form of iron is used to cast articles in foundries such as stoves, pipes, radiators, lamp-posts and rails. Alternatively pig iron may be made into steel (with up to about 2% carbon) or wrought iron (commercially pure iron). Various processes have been used for this, including finery forges, puddling furnaces, Bessemer converters, open hearth furnaces, basic oxygen furnaces, and electric arc furnaces. In all cases, the objective is to oxidize some or all of the carbon, together with other impurities. On the other hand, other metals may be added to make alloy steels.
How Iron was extracted in the 19th century
This heap of iron ore pellets will be used in steel production.
The hardness of the steel depends upon its carbon content, the higher the proportion of carbon, the greater the hardness and the lesser the ductility. The properties of the steel can also be changed by tempering it. To harden the steel, it is heated to red hot and then cooled by quenching it in the water. It becomes harder and more brittle. This steel is then heated to a required temperature and allowed to cool. The steel thus formed is less brittle.
Applications Elemental iron Iron is the most widely used of all the metals, accounting for 95% of worldwide metal production. Its low cost and high strength make it indispensable in engineering applications such as the construction of machinery and machine tools, automobiles, the hulls of large ships, and structural components for buildings. Since pure iron is quite soft, it is most commonly used in the form of steel. Some of the forms in which iron is produced commercially include: • Pig iron has 3.5–4.5% carbon[12] and contains varying amounts of contaminants such as sulfur, silicon and phosphorus. Its only significance is that of an intermediate step on the way from iron ore to cast iron and steel. • Cast iron contains 2–4% carbon, 1–6% silicon, and small amounts of manganese. Contaminants present in pig iron that negatively affect material properties, such as sulfur and phosphorus, have been reduced to an acceptable level. It has a melting point in the range of 1420–1470 K, which is lower than either of its two main components, and
Iron
9 makes it the first product to be melted when carbon and iron are heated together. Its mechanical properties vary greatly, dependent upon the form carbon takes in the alloy. "White" cast irons contain their carbon in the form of cementite, or iron carbide. This hard, brittle compound dominates the mechanical properties of white cast irons, rendering them hard, but unresistant to shock. The broken surface of a white cast iron is full of fine facets of the broken carbide, a very pale, silvery, shiny material, hence the appellation. In grey iron the carbon exists free as fine flakes of graphite, and also renders the material brittle due to the stress-raising nature of the sharp edged flakes of graphite. A newer variant of grey iron, referred to as ductile iron is specially treated with trace amounts of magnesium to alter the shape of graphite to spheroids, or nodules, vastly increasing the toughness and strength of the material.
• Wrought iron contains less than 0.25% carbon.[12] It is a tough, malleable product, but not as fusible as pig iron. If honed to an edge, it loses it quickly. Wrought iron is characterized by the presence of fine fibres of slag entrapped in the metal. Wrought iron is more corrosion resistant than steel. It has been almost completely replaced by mild steel for traditional "wrought iron" products and blacksmithing. Mild steel corrodes more readily that wrought iron, but is cheaper and more widely available. • Carbon steel contains 2.0% carbon or less,[13] with small amounts of manganese, sulfur, phosphorus, and silicon. • Alloy steels contain varying amounts of carbon as
The fining process of smelting iron ore to make wrought iron from pig iron, with the right illustration displaying men working a blast furnace, from the Tiangong Kaiwu encyclopedia, published in 1637 by Song Yingxing.
well as other metals, such as chromium, vanadium, molybdenum, nickel, tungsten, etc. Their alloy content raises their cost, and so they can usually only be justified for specialist uses. Recent developments in ferrous metallurgy have produced a growing range of microalloyed steels, also termed 'HSLA' or high-strength, low alloy steels, containing tiny additions to produce high strengths and often spectacular toughness at minimal cost. The main disadvantage of iron and steel is that pure iron, and most of its alloys, suffer badly from rust if not protected in some way. Painting, galvanization, passivation, plastic coating and bluing are some techniques used to protect iron from rust by excluding water and oxygen or by sacrificial protection.
Iron
10
Iron compounds • Iron oxides (FeO, Fe3O4, and Fe2O3) are ores used for iron production (see bloomery and blast furnace). They are also used as a catalyst in the Space Shuttle Solid Rocket Boosters,[14] and in the production of magnetic storage media in computers. They are often mixed with other compounds, and retain their magnetic properties in solution. • Iron(II) acetate (Fe(CH3CO2)2 (ferrous acetate) is used as a mordant in the dyeing of cloth and leather, and as a wood preservative. • Iron(III) ammonium citrate (C6H5+4yFexNyO7) is used in blueprints.
Iron chloride hexahydrate
• Iron(III) arsenate (FeAsO4) is used in insecticides. • Iron(III) chloride (FeCl3) is used in water purification and sewage treatment, in the dyeing of cloth, as a colouring agent in paints, as an additive in animal feed, and as an etchant for copper in the manufacture of printed circuit boards. • Iron(III) chromate (Fe2(CrO4)3) is a yellow pigment for paints and ceramics. • Iron(III) hydroxide (Fe(OH)3) is used as a brown pigment for rubber and in water purification systems. • Iron(III) phosphate (FePO4) is used in fertilizers and as an additive in human and animal food. • Iron(II) gluconate (Fe(C6H11O7)2) is used as a dietary supplement in iron pills. • Iron(II) oxalate (FeC2O4) is used as yellow pigment for paints, plastics, glass and ceramics, and in photography. • Iron(II) sulfate (FeSO4) is used in water purification and sewage treatment systems, as a catalyst in the production of ammonia, as an ingredient in fertilizer, herbicide, and moss killer, as an additive in animal feed, in wood preservative, and as an additive to flour to increase nutritional iron levels. Experimental iron fertilization of areas of the ocean using iron(II) sulfate has proven successful in increasing plankton growth.[15] [16] [17] The use of iron compounds in organic synthesis is mainly for the reduction of nitro compounds.[18] Additionally, iron has been used for desulfurizations,[19] reduction of aldehydes,[20] and the deoxygenation of amine oxides.[21]
Iron
Biological role Iron is essential to nearly all known organisms. In cells, iron is generally stored in the centre of metalloproteins, because "free" iron (which binds non-specifically to many cellular components) can catalyse production of toxic free radicals. Iron deficiency can lead to iron deficiency anemia. In animals, plants, and fungi, iron is often the metal ion incorporated into the heme complex. Heme is an essential component of cytochrome proteins, which mediate redox reactions, and of oxygen carrier proteins such as hemoglobin, myoglobin, and leghemoglobin. Inorganic iron also contributes to redox reactions in the iron-sulfur clusters of many enzymes, such as Structure of Heme b nitrogenase (involved in the synthesis of ammonia from nitrogen and hydrogen) and hydrogenase. Non-heme iron proteins include the enzymes methane monooxygenase (oxidizes methane to methanol), ribonucleotide reductase (reduces ribose to deoxyribose; DNA biosynthesis), hemerythrins (oxygen transport and fixation in marine invertebrates) and purple acid phosphatase (hydrolysis of phosphate esters). Iron distribution is heavily regulated in mammals, partly because iron has a high potential for biological toxicity[22] . Iron distribution is also regulated because many bacteria require iron, so restricting its availability to bacteria (generally by sequestering it inside cells) can help to prevent or limit infections. This is probably the reason for the relatively low amounts of iron in mammalian milk. A major component of this regulation is the protein transferrin, which binds iron absorbed from the duodenum and carries it in the blood to cells.[23]
Dietary sources Good sources of dietary iron include red meat, fish, poultry, lentils, beans, leaf vegetables, tofu, chickpeas, black-eyed peas, fortified bread, and fortified breakfast cereals. Iron in low amounts is found in molasses, teff and farina. Iron in meat (haem iron) is more easily absorbed than iron in vegetables,[24] but heme/hemoglobin from red meat has effects which may increase the likelihood of colorectal cancer.[25] [26] Iron provided by dietary supplements is often found as iron (II) fumarate, although iron sulfate is cheaper and is absorbed equally well. Elemental iron, despite being absorbed to a much smaller extent (stomach acid is sufficient to convert some of it to ferrous iron), is often added to foods such as breakfast cereals or "enriched" wheat flour (where it is listed as "reduced iron" in the list of ingredients). Iron is most available to the body when chelated to amino acids - iron in this form is ten to fifteen times more bioavailable[27] than any other, and is also available for use as a common iron supplement. Often the amino acid chosen for this purpose is the cheapest and most common amino acid, glycine, leading to "iron glycinate" supplements.[28] The RDA for iron varies considerably based on age, gender, and source of dietary iron (heme-based iron has higher bioavailability).[29] Infants may require iron supplements if they are breast-fed.[30] Blood donors and pregnant women
11
Iron
12
are at special risk of low iron levels and are often advised to supplement their iron intake.
Regulation of uptake Excessive iron can be toxic, because free ferrous iron reacts with peroxides to produce free radicals, which are highly reactive and can damage DNA, proteins, lipids, and other cellular components. Thus, iron toxicity occurs when there is free iron in the cell, which generally occurs when iron levels exceed the capacity of transferrin to bind the iron. Iron uptake is tightly regulated by the human body, which has no regulated physiological means of excreting iron. Only small amounts of iron are lost daily due to mucosal and skin epithelial cell sloughing, so control of iron levels is mostly by regulating uptake.[31] However, large amounts of ingested iron can cause excessive levels of iron in the blood because high iron levels can damage the cells of the gastrointestinal tract, preventing them from regulating iron absorption. The resulting high blood concentrations of iron damage cells in the heart, liver and elsewhere, which can cause serious problems, including long-term organ damage and even death. Humans experience iron toxicity above 20 milligrams of iron for every kilogram of mass, and 60 milligrams per kilogram is a lethal dose.[32] Over-consumption of iron, often the result of children eating large quantities of ferrous sulfate tablets intended for adult consumption, is one of the most common toxicological causes of death in children under six.[32] The DRI lists the Tolerable Upper Intake Level (UL) for adults as 45 mg/day. For children under fourteen years old the UL is 40 mg/day. Regulation of iron uptake is impaired in some people as a result of a genetic defect that maps to the HLA-H gene region on chromosome 6. In these people, excessive iron intake can result in iron overload disorders, such as hemochromatosis. Many people have a genetic susceptibility to iron overload without realizing it or being aware of a family history of the problem. For this reason, it is advised that people do not take iron supplements unless they suffer from iron deficiency and have consulted a doctor. Hemochromatosis is estimated to cause disease in between 0.3 and 0.8% of Caucasians.[33] The medical management of iron toxicity is complex, and can include use of a specific chelating agent called deferoxamine to bind and expel excess iron from the body.[34] .
See also • • • • • • • • •
El Mutún in Bolivia, where 20% of the world's accessible iron and magnesium is located. Iron Age Iron fertilization - Proposed fertilization of oceans to stimulate phytoplankton growth. Iron (metaphor) Iron in mythology List of countries by iron production Pelletising - Process of creation of iron ore pellets. Rustproof iron Specht Building - A historic landmark in Omaha, Nebraska utilizing an iron facade.
Iron
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Books • Doulias PT, Christoforidis S, Brunk UT, Galaris D. Endosomal and lysosomal effects of desferrioxamine: protection of HeLa cells from hydrogen peroxide-induced DNA damage and induction of cell-cycle arrest. Free Radic Biol Med. 2003;35:719-28. • H. R. Schubert, History of the British Iron and Steel Industry ... to 1775 AD (Routledge, London, 1957) • R. F. Tylecote, History of Metallurgy (Institute of Materials, London 1992). • R. F. Tylecote, 'Iron in the Industrial Revolution' in J. Day and R. F. Tylecote, The Industrial Revolution in Metals (Institute of Materials 1991), 200-60.
External links • • • •
WebElements.com – Iron [35] It's Elemental – Iron [36] The Most Tightly Bound Nuclei Crystal structure of iron [38]
[37]
References [1] " Iron and Nickel Abundances in H~II Regions and Supernova Remnants (http:/ / www. aas. org/ publications/ baas/ v27n2/ aas186/ abs/ S3707. html)". June 14, 1995. . Retrieved 2008-05-21.. [2] " Hardness Conversion Chart (http:/ / mdmetric. com/ tech/ hardnessconversion. html)". . Retrieved 2009-07-07. [3] V. Raghavan (2004). Materials Science and Engineering (http:/ / books. google. com/ books?id=LgB5dkmPML0C& pg=PA218). PHI Learning Pvt. Ltd.. p. 218. ISBN 8120324552. . [4] " Properties of Various Pure Irons : Study on pure iron I (http:/ / ci. nii. ac. jp/ naid/ 110001459778/ en)". Tetsu-to-Hagane 50 (1): 42–47. . [5] John Wilson Martin (2007). Concise encyclopedia of the structure of materials (http:/ / books. google. com/ books?id=xv420pEC2qMC& pg=PA183). Elsevier. p. 183. ISBN 0080451276. . [6] " Iron: geological information (http:/ / www. webelements. com/ iron/ geology. html)". . Retrieved 2008-05-21.. [7] Dauphas, N. & Rouxel, O. 2006. Mass spectrometry and natural variations of iron isotopes. Mass Spectrometry Reviews, 25, 515-550 [8] The origins of Iron Working in India: New evidence from the Central Ganga plain and the Eastern Vindhyas by Rakesh Tewari (Director, U.P. State Archaeological Department) (http:/ / antiquity. ac. uk/ projgall/ tewari/ tewari. pdf) [9] E. Photos, 'The Question of Meteoritic versus Smelted Nickel-Rich Iron: Archaeological Evidence and Experimental Results' World Archaeology Vol. 20, No. 3, Archaeometallurgy (Feb., 1989), pp. 403-421. [10] Muhly, James D. 'Metalworking/Mining in the Levant' pp. 174-83 in Near Eastern Archaeology ed. S. Richard Winona Lake, IN: Eisenbrauns (2003): 180. [11] Donald B. Wagner, 'Chinese blast furnaces from the 10th to the 14th century' Historical Metallurgy 37(1) (2003), 25-37; originally published in West Asian Science, Technology, and Medicine 18 (2001), 41-74. [12] Camp, James McIntyre; Francis, Charles Blaine (1920). The Making, Shaping and Treating of Steel (http:/ / books. google. com/ books?id=P9MxAAAAMAAJ). Pittsburgh: Carnegie Steel Company. pp. 173–174. . [13] " Classification of Carbon and Low-Alloy Steels (http:/ / www. key-to-steel. com/ Articles/ Art62. htm)". . Retrieved 2008-01-05. [14] http:/ / science. ksc. nasa. gov/ shuttle/ technology/ sts-newsref/ srb. html [15] Vivian Marx (2002). " The Little Plankton That Could…Maybe (http:/ / www. sciam. com/ article. cfm?articleID=000A5750-8AC2-1D9C-815A809EC5880000)". Scientific American. . [16] Melinda Ferguson, David Labiak, Andrew Madden, Joseph Peltier. " The Effect of Iron on Plankton Use of CO2 (http:/ / www. cem. msu. edu/ ~cem181h/ projects/ 96/ iron/ cem. html)". CEM 181H. . Retrieved 2007-05-05. [17] Dopyera, Caroline (October 1996). " The Iron Hypothesis (http:/ / www. palomar. edu/ oceanography/ iron. htm)". EARTH. . Retrieved 2007-05-05. [18] Fox, B. A.; Threlfall, T. L. (1973). Organic Syntheses 5: 346.; Vol. 44, p.34 (1964). ( Article (http:/ / www. orgsyn. org/ orgsyn/ prep. asp?prep=cv5p0346))
Iron
14
[19] Blomquist, A. T.; Dinguid, L. I. (1947). J. Org. Chem. 12&page=718. [20] Clarke, H. T.; Dreger, E. E. (1941). Organic Synthesis 1: 304.; Vol. 6, p.52 (1926). ( Article (http:/ / www. orgsyn. org/ orgsyn/ prep. asp?prep=cv1p0304)). [21] den Hertog, J.; Overhoff (1950). Recl. Trav. Chim. 69: 468. [22] http:/ / cat. inist. fr/ ?aModele=afficheN& cpsidt=17328512 [23] Tracey A. Rouault. " How Mammals Acquire and Distribute Iron Needed for Oxygen-Based Metabolism (http:/ / biology. plosjournals. org/ perlserv/ ?request=get-document& doi=10. 1371/ journal. pbio. 0000079)". . Retrieved 2006-06-19. [24] Food Standards Agency - Eat well, be well - Iron deficiency (http:/ / www. eatwell. gov. uk/ healthissues/ irondeficiency/ ) [25] Sesink AL, Termont DS, Kleibeuker JH, Van der Meer R (1999). " Red meat and colon cancer: the cytotoxic and hyperproliferative effects of dietary heme (http:/ / cancerres. aacrjournals. org/ cgi/ pmidlookup?view=long& pmid=10582688)". Cancer Research (http:/ / cancerres. aacrjournals. org/ ) 59 (22): 5704–9. PMID 10582688. . [26] Glei M, Klenow S, Sauer J, Wegewitz U, Richter K, Pool-Zobel BL (2006). "Hemoglobin and hemin induce DNA damage in human colon tumor cells HT29 clone 19A and in primary human colonocytes". Mutat. Res. 594 (1-2): 162–71. doi: 10.1016/j.mrfmmm.2005.08.006 (http:/ / dx. doi. org/ 10. 1016/ j. mrfmmm. 2005. 08. 006). PMID 16226281. [27] Pineda O, Ashmead HD (2001). "Effectiveness of treatment of iron-deficiency anemia in infants and young children with ferrous bis-glycinate chelate". Nutrition 17 (5): 381–4. doi: 10.1016/S0899-9007(01)00519-6 (http:/ / dx. doi. org/ 10. 1016/ S0899-9007(01)00519-6). PMID 11377130. [28] Ashmead, H. DeWayne (1989). Conversations on Chelation and Mineral Nutrition. Keats Publishing. ISBN 0-87983-501-X. [29] " Dietary Reference Intakes: Elements (http:/ / www. iom. edu/ Object. File/ Master/ 7/ 294/ 0. pdf)" (PDF). The National Academies. 2001. . Retrieved 2008-05-21. [30] " Iron Deficiency Anemia (http:/ / bodyandhealth. canada. com/ condition_info_details. asp?disease_id=274)" (web page). MediResource. . Retrieved 2008-12-17. [31] Kumar, Vinay; Abbas, Abul K; Fausto, Nelson (2005). " Anemia (http:/ / www. mdconsult. com/ das/ book/ body/ 90234262-4/ 0/ 1249/ 121. html)". Robbins and Cotran: Pathologic Basis of Disease, 7th edition. Elsevier Saunders. . Retrieved 2008-03-14. [32] " Toxicity, Iron (http:/ / www. emedicine. com/ emerg/ topic285. htm)". Emedicine. . Retrieved 2006-06-19. [33] Durupt S, Durieu I, Nove-Josserand R, et al.: [Hereditary hemochromatosis]. Rev Med Interne 2000 Nov; 21(11): 961-71[Medline]. [34] Miller, Marvin J. (1989-11-01). "Syntheses and therapeutic potential of hydroxamic acid based siderophores and analogs". Chemical Reviews 89 (7): 1563–1579. doi: 10.1021/cr00097a011 (http:/ / dx. doi. org/ 10. 1021/ cr00097a011). [35] [36] [37] [38]
http:/ / www. webelements. com/ webelements/ elements/ text/ Fe/ index. html http:/ / education. jlab. org/ itselemental/ ele026. html http:/ / hyperphysics. phy-astr. gsu. edu/ hbase/ nucene/ nucbin2. html http:/ / www. webelements. com/ webelements/ elements/ text/ Fe/ xtal. html
Article Sources and Contributors
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15
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16
Cobalt
1
Cobalt iron ← cobalt → nickel ↑ Co ↓ Rh
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
27Co Periodic table
Appearance hard lustrous gray metal
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties ColorDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
1790
1960
2165
2423
2755
3198
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1648 kJ·mol−1 3rd: 3232 kJ·mol−1Atomic radiusCovalent radius Miscellaneous
Cobalt
2
Crystal structureMagnetic orderingElectrical resistivityThermal conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusPoisson ratioMohs hardnessVickers hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of cobalt iso
N.A.
half-life
DM
DE (MeV)
DP
56
syn
77.27 d
ε
4.566
56
57
syn
271.79 d
ε
0.836
57
58
syn
70.86 d
ε
2.307
58
59
100%
59
60
syn
5.2714 years
2.824
60
Co Co Co Co Co
Fe Fe Fe
Co is stable with 32 neutron β−,γ,γ
Ni
cobalt, Co, 27 transition metal9, 4, d58.933195(5) g·mol−1 [Ar] 4s2 3d7 2, 8, 15, 2 (Image) metallic gray 8.90 g·cm−3 7.75 g·cm−3 1768 K,1495 °C,2723 °F 3200 K,2927 °C,5301 °F 16.06 kJ·mol−1 377 kJ·mol−1 (25 °C) 24.81 J·mol−1·K−1 5, 4 , 3, 2, 1 [1] , -1 (amphoteric oxide) 1.88 (Pauling scale) 1st: 760.4 kJ·mol−1125 pm126±3 (low spin), 150±7 (high spin) pm hexagonal ferromagnetic (20 °C) 62.4 nΩ·m (300 K) 100 W·m−1·K−1 (25 °C) 13.0 µm·m−1·K−1 (20 °C) 4720 m/s 209 GPa 75 GPa 180 GPa 0.31 5.0 1043 MPa 700 MPa 7440-48-4 Cobalt (pronounced /'kəʊbɒlt/)[2] is a hard, lustrous, grey metal, a chemical element with symbol Co and atomic number 27. Although cobalt-based colors and pigments have been used since ancient times for making jewelry and paints, and miners have long used the name kobold ore for some minerals, the free metallic cobalt was not prepared and discovered until 1735 by Georg Brandt. Cobalt is found in various metallic-lustered ores for example cobaltite (CoAsS), but it is produced as a by-product of copper and nickel mining. The copper belt in the Democratic Republic of the Congo and Zambia yields most of the worldwide mined cobalt. Cobalt is used in the preparation of magnetic, wear-resistant, and high-strength alloys. Cobalt blue (cobalt(II) aluminate, CoAl2O4) gives a distinctive deep blue color to glass, ceramics, inks, paints, and varnishes. Cobalt-60 is a commercially important radioisotope, used as a tracer and in the production of gamma rays for industrial use. Cobalt is an essential trace-element for all multicellular organisms as the active center of coenzymes called cobalamins. These include vitamin B-12 which is essential for mammals. Cobalt is also an active nutrient for bacteria, algae, and fungi, and may be a necessary nutrient for all life.
Characteristics Cobalt is a ferromagnetic metal. Pure cobalt is not found in nature, but compounds of cobalt are common. Small amounts of it are found in most rocks, soil, plants, and animals. It is the element of atomic number 27. The Curie temperature is 1115 °C, and the magnetic moment is 1.6–1.7 Bohr magnetons per atom. In nature, it is frequently associated with nickel, and both are characteristic minor components of meteoric iron. Mammals require small amounts of cobalt which is the basis of vitamin B12. Cobalt-60, an artificially produced radioactive isotope of cobalt, is an important radioactive tracer and cancer-treatment
Cobalt
3
agent. Cobalt has a relative permeability two thirds that of iron. Metallic cobalt occurs as two crystallographic structures: hcp and fcc. The ideal transition temperature between hcp and fcc structures is 450 °C, but in practice, the energy difference is so small that random intergrowth of the two is common.[3]
History Cobalt compounds have been used for centuries to impart a rich blue color to glass, glazes, and ceramics. Cobalt has been detected in Egyptian sculpture and Persian jewelry from the third millennium BC, in the ruins of Pompeii (destroyed AD 79), and in China dating from the Tang dynasty (AD 618–907) and the Ming dynasty (AD 1368–1644)[4] . Cobalt glass ingots have been recovered from the Uluburun shipwreck, dating to the late 14th century BC.[5] Swedish chemist Georg Brandt (1694–1768) is credited with isolating cobalt circa 1735.[6] He was able to show that cobalt was the source of the blue color in glass, which previously had been attributed to the bismuth found with cobalt. The word cobalt is derived from the German kobalt, from kobold meaning "goblin", a term used for the ore of cobalt by miners. The first attempts at smelting the cobalt ores to produce cobalt metal failed, yielding cobalt(II) oxide instead. Also, because the primary ores of cobalt always contain arsenic, smelting the ore oxidized into the highly toxic and volatile oxide As4O6, which was inhaled by workers. During the 19th century, cobalt blue was produced at the Norwegian Blaafarveværket (70–80% of world production), led by the Prussian industrialist Benjamin Wegner. In 1938, John Livingood and Glenn Seaborg discovered cobalt-60. This isotope was famously used at Columbia University in the 1950s to establish parity violation in beta decay.
Occurrence Cobalt occurs in copper and nickel minerals and in combination with sulfur and arsenic in the sulfidic cobaltite (CoAsS), safflorite (CoAs2) and skutterudite (CoAs3) minerals.[7] The mineral cattierite is similar to pyrite and occurs together vaesite in the copper deposits in the Katanga Province.[8] If the sulfides come in contact with the atmosphere weathering starts transforming the minerals by oxidation. The products of the oxidation are for example pink erythrite ('cobalt glance': Co3(AsO4)2·8H2O) and sphaerocobaltite (CoCO3).
Production Cobalt is not found as a native metal but generally found in the form of ores. Cobalt is usually not mined alone, and tends to be produced as a by-product of nickel and copper mining activities. The main ores of cobalt are cobaltite, erythrite, glaucodot, and skutterudite.[9] [10] In 2005, the copper deposits in the Katanga Province (former Shaba province) of the Democratic Republic of Cobalt ore
Cobalt
the Congo was the top producer of cobalt with almost 40% world share, reports the British Geological Survey.[11] The problematic political situation in the Congo influences the price of cobalt significantly, best example was the Shaba crisis in 1978.[12] Cobalt output in 2005 There are several methods which can be used to separate cobalt from copper and nickel. They depend on the concentration of cobalt and the exact composition of the used ore. The first possible separation step is the froth flotation of the ore, in which special surfactants yield in an enrichment of cobalt. The following roasting of the ores can be conducted in a way that the cobalt sulfide is oxidized to the cobalt sulfate, while the copper and the iron are oxidized to the oxide. The leaching with water extracts the sulphate together with the arsenates. The residues are World production trend further leached with sulfuric acid yielding a solution of copper sulfate. They also present iron nickel and cobalt salts can be precipitated by chlorine or hypochloride. If the copper is not produced by leaching and electrowinning but by the pyrometallurgic process, the cobalt can be leached from the slag of the copper smelter.[13]
All the above mentioned processes yield copper compounds which are transformed into the cobalt oxide Co3O4. The reduction to the metal is done either by the aluminothermic reaction or reduction with carbon in a blast furnace.[7] In 2008, The London Metal Exchange announced that Cobalt would be be traded as a commodity on the London Metal Exchange.[14]
Isotopes 59
Cobalt is the only stable cobalt isotope. 22 radioisotopes have been characterized with the most stable being 60Co with a half-life of 5.2714 years, 57Co with a half-life of 271.79 days, 56Co with a half-life of 77.27 days, and 58Co with a half-life of 70.86 days. All of the remaining radioactive isotopes have half-lives that are less than 18 hours, and the majority of these are less than 1 second. This element also has 4 meta states, all of which have half-lives less than 15 minutes. The isotopes of cobalt range in atomic weight from 50 u (50Co) to 73 u (73Co). The primary decay mode for isotopes with atomic mass unit values less than that of the most abundant stable isotope, 59Co, is electron capture and the primary mode of decay for those of greater than 59 atomic mass units is beta decay. The primary decay products before 59Co are element 26 (iron) isotopes and the primary products after are element 28 (nickel) isotopes.
4
Cobalt Cobalt radioisotopes in medicine Cobalt-60 (Co-60 or 60Co) is a radioactive metal that is used in radiotherapy. It produces two gamma rays with energies of 1.17 MeV and 1.33 MeV. The 60Co source is about 2 cm in diameter and as a result produces a geometric penumbra, making the edge of the radiation field fuzzy. The metal has the unfortunate habit of producing a fine dust, causing problems with radiation protection. Cobalt-60 has a radioactive half-life of 5.27 years. This decrease in activity requires periodic replacement of the sources used in radiotherapy and is one reason why cobalt machines have been largely replaced by linear accelerators in modern radiation therapy. Cobalt from radiotherapy machines has been a serious hazard when not disposed of properly, and one of the worst radiation contamination accidents in North America occurred in 1984, after a discarded cobalt-60 containing radiotherapy unit was mistakenly disassembled in a junkyard in Juarez, Mexico.[15] Cobalt-57 (Co-57 or 57Co) is a cobalt radioisotope most often used in medical tests, as a radiolabel for vitamin B12 uptake, and for the Schilling test.[16]
Industrial uses for radioactive isotopes Cobalt-60 (Co-60 or 60Co) is useful as a gamma ray source because it can be produced in predictable quantity and high activity by simply exposing natural cobalt to neutrons in a reactor for a period. Its uses include sterilization of medical supplies and medical waste, radiation treatment of foods for sterilization (cold pasteurization), industrial radiography (e.g., weld integrity radiographs), density measurements (e.g., concrete density measurements), and tank fill height switches. Cobalt-57 is used as a source in Mössbauer spectroscopy and is one of several possible sources in XRF devices (Lead Paint Spectrum Analyzers).
Cobalt-60 as weapon Nuclear weapon designs could intentionally incorporate 59Co, some of which would be activated in a nuclear explosion to produce 60Co. The 60Co, dispersed as nuclear fallout, creates what is sometimes called a dirty bomb or cobalt bomb.[17]
Applications Alloys The cobalt based superalloys consume most of the produced cobalt. The temperature stability of these alloys makes them suitable for turbine blades within gas turbines and jet aircraft engines. The nickel-based single crystal alloys surpass the cobalt based in temperature stability, but the cobalt based are still in use. These alloys are also corrosion and wear-resistant.[18] Special cobalt chromium molybdenum alloys are used for prosthetic parts such as hip and knee replacements.[19] Cobalt alloys are also used for dental prosthetics, where they are useful to avoid allergies to nickel.[20] Some high speed steels also use cobalt to increase heat and wear-resistance. The special alloys of aluminium, nickel, cobalt and iron, known as Alnico, and of samarium and cobalt (samarium-cobalt magnet) are used in permanent magnets.[21]
5
Cobalt
6
Batteries Lithium cobalt oxide (LiCoO2) is widely used in Lithium ion battery electrodes.[22] Nickel-cadmium (NiCd) and nickel metal hydride (NiMH) batteries also contain significant amounts of cobalt.
Catalyst Several cobalt compounds are used in chemical reactions as catalysts. Cobalt acetate is used for the production of terephthalic acid as well as dimethyl terephthalic acid, which are key compounds in the production of Polyethylene terephthalate. The steam reforming and hydrodesulfuration for the production of petroleum, which uses mixed cobalt molybdenum aluminium oxides as a catalyst, is another important application.[22] Cobalt and its compounds, especially cobalt carboxylates (known as cobalt soaps), are good oxidation catalysts. They are used in paints, varnishes, and inks as drying agents through the oxidation of certain compounds.[22] The same carboxylates are used to improve the adhesion of the steel to rubber in steel-belted radial tires.[22]
Pigments and coloring Before the 19th century, the predominant use of cobalt was the pigmentation of glass. The colors cobalt blue and cobalt green originated from this use.[23] Cobalt has been used to color glass since the Bronze Age. The excavation of the Uluburun shipwreck yielded an ingot of blue glass which was cast during the 14th century BC.[24] Blue glass items from Egypt are colored with copper, iron, or cobalt. The oldest cobalt-colored glass was from the time of the Eighteenth dynasty (1550–1292 BC). The location where the cobalt compounds were obtained is unknown.[25] [26]
Cobalt blue glass
Other uses • Electroplating due to its appearance, hardness, and resistance to oxidation • Ground coats for porcelain enamels • Purification of histidine-tagged fusion proteins in biotechnology applications
Compounds Common oxidation states of cobalt include +2 and +3, although compounds with oxidation state +1 are also known. The most stable oxidation state for simple compounds is +2. Cobalt(II) salts form the red-pink [Co(H2O)6]2+ complex in aqueous solution. Adding excess chloride will change the color from pink to blue, due to the formation of [CoCl4]2-.
Cobalt
Chalcogen compounds Several oxides of cobalt are known. Green cobalt(II) oxide (CoO) has NaCl structure and is readily oxidized with water and oxygen to brown cobalt(III) hydroxide (Co(OH)3). At temperatures of 400–500 °C the CoO is oxidized to the blue cobalt(II,III) oxide (Co3O4), which has spinel structure. The brown cobalt(III) oxide (Co2O3) is the least stable of the oxides. Cobalt oxides are antiferromagnetic at low temperature: CoO (Neel temperature 291 K) and Co3O4 (Neel temperature: 40 K), which is analogous to magnetite (Fe3O4), with a mixture of +2 and +3 oxidation states. The oxide Co2O3 is probably unstable; it has never been synthesized. The sulfur compounds are the two black cobalt(II) sulfide (CoS2) and cobalt(III) sulfide (Co2S3).
Halogen compounds The halogen compounds of cobalt are cobalt(II) fluoride (CoF2), cobalt(II) chloride (CoCl2), cobalt(II) bromide (CoBr2), cobalt(II) iodide (CoI2), and cobalt(III) fluoride (CoF3). Cobalt(II) chloride is commonly found as an indicator of dryness in silica gel beads used as a desiccant. Anhydrous cobalt(II) chloride is blue, while the hexahydrate is red. The reduction potential for Co3+ + e− → Co2+ (+1.92 V) is far beyond the one for chlorine. As a consequence, only the fluoride is able to exist in the cobalti-status.
Coordination compounds Other than Co3O4 and the brown fluoride CoF3 (which is instantly hydrolyzed in water), all compounds containing cobalt in the +3 oxidation state are stabilized by complex ion formation. Examples for the more exotic oxidation states +1, +4 and +5 are the compounds tris(triphenylphosphine)cobalt(I) chloride ((P(C6H5)3)3CoCl), caesium hexafluorocobaltate (Cs2CoF6)) and potassium percobaltate (K3CoO4).[7] The class of vitamin B12 compounds are coordination complexes of elaborated corrin rings with a central cobalt atom. Alfred Werner, a pioneer in coordination chemistry, worked with compounds of empirical formula CoCl3(NH3)6; one of the isomers determined was cobalt(III) hexammine chloride. This coordination complex, a "typical" Werner-type complex, consists of a central cobalt atom coordinated by six ammine ligands orthogonal to each other, and three chloride counteranions. Using chelating ethylenediamine ligands in place of ammonia gives tris(ethylenediamine)cobalt(III) chloride ([Co(en)3]Cl), which was one of the first coordination complex showing stereochemistry. The complex can take either right- or left-handed forms of a three-bladed propeller. This complex was first isolated by Werner as yellow-gold needle-like crystals.[27] Cobaltocene is a fairly stable cobalt analog to ferrocene.
7
Cobalt
8
Biological role Cobalt in small amounts is essential to many living organisms, including humans. Having 0.13 to 0.30 mg/kg of cobalt in soils markedly improves the health of grazing animals. Cobalt is a central component of the vitamin cobalamin, or vitamin B12. Although cobalt proteins are less common than proteins containing metals like manganese, iron, or zinc, several are known. Most cobalt proteins use a cofactor based on the corrin cobalt, derived from vitamin B12, but there are also a few proteins known in which cobalt is directly coordinated by the protein structure; Methionine aminopeptidase 2 and Nitrile hydratase are two examples.[28]
Precautions Although cobalt is an essential element for life in
Cobalamin
minute amounts, at higher levels of exposure it shows mutagenic and carcinogenic effects similar to nickel (see Cobalt Poisoning).[29] In 1966, the addition of cobalt compounds to stabilize beer foam in Canada led to cardiomyopathy, which came to be known as beer drinker's cardiomyopathy.[30] Powdered cobalt in metal form is a fire hazard. After nickel and chromium, cobalt is a major cause of contact dermatitis.[31]
External links • National Pollutant Inventory - Cobalt fact sheet • WebElements.com – Cobalt [33]
[32]
• London celebrates 50 years of Cobalt-60 Radiotherapy
[34]
References [1] Choi, Hyungsoo (2003). "Liquid Cobalt (I) Hydride Complexes as Precursors for Chemical Vapor Deposition". Chemistry of Materials 15: 3121. doi: 10.1021/cm030100e (http:/ / dx. doi. org/ 10. 1021/ cm030100e). [2] [3] [4] [5]
Oxford English Dictionary, 2nd Edition 1989. " Properties and Facts for Cobalt (http:/ / www. americanelements. com/ co. html)". . Retrieved 2008-09-19. Encyclopedia Britannica Online. Pulak, Cemal (1998). "The Uluburun shipwreck: an overview". International Journal of Nautical Archaeology 27 (3): 188–224. doi: 10.1111/j.1095-9270.1998.tb00803.x (http:/ / dx. doi. org/ 10. 1111/ j. 1095-9270. 1998. tb00803. x). [6] Wang, Shijie (2006). "Cobalt—Its recovery, recycling, and application". Journal of the Minerals, Metals and Materials Society 58 (10): 47–50. doi: 10.1007/s11837-006-0201-y (http:/ / dx. doi. org/ 10. 1007/ s11837-006-0201-y). [7] Holleman, A. F., Wiberg, E., Wiberg, N. (2007). "Cobalt" (in German). Lehrbuch der Anorganischen Chemie, 102nd ed.. de Gruyter. pp. 1146–1152. ISBN 9783110177701. [8] Kerr, Paul F. (1945). Cattierite and Vaesite: New Co-Ni Minerals from the Belgian Kongo (http:/ / www. minsocam. org/ ammin/ AM30/ AM30_483. pdf). 30. pp. 483–492. . [9] Shedd, Kim B.. " Mineral Yearbook 2006: Cobalt (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ cobalt/ myb1-2006-cobal. pdf)". United States Geological Survey. . Retrieved 2008-10-26.
Cobalt [10] Shedd, Kim B.. " Commodity Report 2008: Cobalt (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ cobalt/ mcs-2008-cobal. pdf)". United States Geological Survey. . Retrieved 2008-10-26. [11] " African Mineral Production (http:/ / www. bgs. ac. uk/ mineralsuk/ downloads/ african_mp_01_05. pdf)". British Geological Survey. . Retrieved 2009-06-06. [12] Wellmer, Friedrich-Wilhelm. " Global Nonfuel Mineral Resources and Sustainability (http:/ / pubs. usgs. gov/ circ/ 2007/ 1294/ paper1. html)". . Retrieved 2009-05-16. [13] Joseph R. Davis (2000). ASM specialty handbook: nickel, cobalt, and their alloys (http:/ / books. google. com/ books?id=IePhmnbmRWkC& dq=cobalt+ copper+ nickel+ ore+ separate& lr=& num=100& as_brr=3& source=gbs_navlinks_s). ASM International. p. 347. ISBN 0871706857. . [14] " LME to launch minor metals contracts in H2 2009 (http:/ / lme. com/ 6241. asp)". London Metal Exchange. 4 September 2009. . Retrieved 28 July 2009. [15] " The Juarez accident (http:/ / query. nytimes. com/ gst/ fullpage. html?sec=health& res=9501E7D71338F932A35756C0A962948260)". New York Times. . Retrieved 2009-06-06. [16] " An overview of cobalt radioisotopes in medicine (http:/ / www. helium. com/ items/ 876792-an-overview-of-cobalt-radioisotopes-in-medicine)". . Retrieved 2009-06-06. [17] Payne, L.R. (1977). " The Hazards of Cobalt (http:/ / occmed. oxfordjournals. org/ cgi/ content/ abstract/ 27/ 1/ 20)". Occupational Medicine 27: 20–25. doi: 10.1093/occmed/27.1.20 (http:/ / dx. doi. org/ 10. 1093/ occmed/ 27. 1. 20). . [18] Donachie, Matthew J. (2002). Superalloys: A Technical Guide (http:/ / books. google. de/ books?id=vjCJ5pI1QpkC). ASM International. ISBN 9780871707499. . [19] Michel, R.; Nolte, M.; Reich M.; Löer, F. (1991). "Systemic effects of implanted prostheses made of cobalt-chromium alloys". Archives of Orthopaedic and Trauma Surgery 110 (2): 61–74. doi: 10.1007/BF00393876 (http:/ / dx. doi. org/ 10. 1007/ BF00393876). [20] Disegi, John A. (1999). Cobalt-base Aloys for Biomedical Applications (http:/ / books. google. com/ books?id=z4rXM1EnPugC). ASTM International. ISBN 0803126085. ., p 34 [21] Luborsky, F. E.; Mendelsohn, L. I.; Paine, T. O. (1957). "Reproducing the Properties of Alnico Permanent Magnet Alloys with Elongated Single-Domain Cobalt-Iron Particles". Journal Applied Physics 28 (344): 344. doi: 10.1063/1.1722744 (http:/ / dx. doi. org/ 10. 1063/ 1. 1722744). [22] Hawkins, M. (2001). "Why we need cobalt". Applied Earth Science: Transactions of the Institution of Mining & Metallurgy, Section B 110 (2): 66–71. [23] Venetskii, S. (1970). "The charge of the guns of peace". Metallurgist 14 (5): 334–336. doi: 10.1007/BF00739447 (http:/ / dx. doi. org/ 10. 1007/ BF00739447). [24] Henderson, Julian (2000). " Glass (http:/ / books. google. com/ books?id=p9xJ-VpUuNkC)". The Science and Archaeology of Materials: An Investigation of Inorganic Materials. Routledge. p. 60. ISBN 9780415199339. . [25] Rehren, Th. (2003). "Aspects of the Production of Cobalt-blue Glass in Egypt". Archaeometry 43 (4): 483–489. doi: 10.1111/1475-4754.00031 (http:/ / dx. doi. org/ 10. 1111/ 1475-4754. 00031). [26] Lucas, A. (2003). Ancient Egyptian Materials and Industries (http:/ / books. google. com/ books?id=GugkliLHDMoC). Kessinger Publishing. p. 217. ISBN 9780766151413. . [27] A. Werner (1912). "Zur Kenntnis des asymmetrischen Kobaltatoms. V". Chemische Berichte 45: 121–130. doi: 10.1002/cber.19120450116 (http:/ / dx. doi. org/ 10. 1002/ cber. 19120450116). [28] Kobayashi, Michihiko; Shimizu, Sakayu (1999). "Cobalt proteins". European Journal of Biochemistry 261 (1): 1–9. doi: 10.1046/j.1432-1327.1999.00186.x (http:/ / dx. doi. org/ 10. 1046/ j. 1432-1327. 1999. 00186. x). [29] " Report on Carcinogens, Eleventh Edition: Cobalt Sulfate (http:/ / ntp. niehs. nih. gov/ ntp/ roc/ eleventh/ profiles/ s048zcob. pdf)". National Toxicology Program. . Retrieved 2008-11-13. [30] Donald G. Barceloux; Donald Barceloux (1999). "Cobalt". Clinical Toxicology 37 (2): 201–216. doi: 10.1081/CLT-100102420 (http:/ / dx. doi. org/ 10. 1081/ CLT-100102420). [31] Basketter, David A.; Angelini, Gianni; Ingber, Arieh; Kern, Petra S.; Menné, Torkil (2003). "Nickel, chromium and cobalt in consumer products: revisiting safe levels in the new millennium". Contact Dermatitis 49 (1): 1–7. doi: 10.1111/j.0105-1873.2003.00149.x (http:/ / dx. doi. org/ 10. 1111/ j. 0105-1873. 2003. 00149. x). [32] http:/ / www. npi. gov. au/ database/ substance-info/ profiles/ 26. html [33] http:/ / www. webelements. com/ webelements/ elements/ text/ Co/ key. html [34] http:/ / www. caro-acro. ca/ caro/ educ/ publ/ vig/ vignettes/ cobalt/ Interactions. pdf
9
Article Sources and Contributors
Article Sources and Contributors Cobalt Source: http://en.wikipedia.org/w/index.php?oldid=304835490 Contributors: *drew, 128.139.197.xxx, 24.15.135.xxx, Addshore, Ahoerstemeier, Aitias, Alansohn, Alba, Aldaniel, AlexiusHoratius, Allstarecho, Alon, Andres, AndrewDMcGregor, Andrewrp, Andy M. Wang, Antandrus, Anwar saadat, Ap, Apple2, Apyule, Arcadian, Archimerged, Arthurzahn, Aussie Alchemist, Baccyak4H, Bbi5291, Bcorr, Beetstra, Bender235, Beneatsfood, Benjiboi, Bert.Kilanowski, Bhadani, BiggKwell, Bilbobee, BlueEarth, Bobo192, Bodnotbod, Bogey97, Bokunenjin, Bomac, Borislav.dopudja, BrandonKloppenburgPh.D, Brian Huffman, Bryan Derksen, Bwbrian, Cadmium, Caesura, CambridgeBayWeather, Can't sleep, clown will eat me, Capricorn42, CardinalDan, Carnildo, Celarnor, Centrx, CeresVesta, CesarAndreu, Chameleon, Chiu frederick, Chris 73, Chris Roy, Chrislk02, Christopher Thomas, Circeus, ClintReuter, Condem, Conversion script, Coppertwig, Cornthwaite, Corpx, Crystallina, Cst17, Cyrius, Daa89563, Daduzi, Daibhid C, Daniel5127, DanielDeibler, Darrien, Darth Panda, David Latapie, Dawn Bard, DeadEyeArrow, Deli nk, Delirium, Delta G, DerHexer, Discospinster, DnL7up, Donarreiskoffer, Doulos Christos, Drivenapart, Dwmyers, Dwtrojans, Edgar181, Eedo Bee, Eeekster, Eigenlambda, El C, Eldin raigmore, Elkman, Emhoo, Emperorbma, Epbr123, Fatal!ty, Femto, Frankenpuppy, Fredrik, Friginator, Funny Muffin, Fvw, Gene Nygaard, Giftlite, Gilgamesh, Giqi, Glacier Wolf, Gman124, Goldom, GraemeL, Graibeard, Grendelkhan, Griffin7890, Gscshoyru, Gurch, Gwernol, Hak-kâ-ngìn, HappyM, Hdt83, Hede2000, Heimstern, Hephaestos, HereToHelp, Heygogirl001, I HATE STARWARS, IW.HG, IainP, Icairns, Iepeulas, Imjdk, Insanity Incarnate, Ioscius, Iridescent, Irishguy, Itai, J.delanoy, JForget, JPINFV, Jacek Kendysz, Jacob.jose, Janus Shadowsong, Jaraalbe, Jarsyl, JayJasper, JeremyA, Jerzy, Jgrimmer, Jiang, Jj137, Joanjoc, JoanneB, John, John254, Jojhutton, Jonpro, Jose77, Josephgrossberg, Josh2Go, KT322, Karl-Henner, Karlhahn, Katieh5584, Kazvorpal, Kbdank71, Keilana, Kelleyrocks2008, Kingpin13, KnowledgeOfSelf, Koavf, Kpjas, Kreachure, Kubigula, Kurykh, Kwamikagami, Kwsn, LA2, LMB, LarryMorseDCOhio, LeaveSleaves, LeeG, Leonard^Bloom, Lethalgeek, Leyo, LibLord, Lil blak sam, Luk, MER-C, MZMcBride, Mad apothecary, Maddie!, Madris, Mani1, Marc Tobias Wenzel, Materialscientist, Mav, Mclowes, Mderezynski, Mdf, Meekywiki, Mejor Los Indios, Mentifisto, Mgimpel, Mifter, Mike.rooke, Mild Bill Hiccup, Minesweeper, Mmccrae, Mr. Lefty, N2e, Nabla, Naughtydog05, Nergaal, Nick Y., Nico, Nihiltres, Nikolabc, Nimur, NobleHelium, NoneTooSoon, Oda Mari, Oliver Lineham, Oxymoron83, PJM, Pakaran, Persian Poet Gal, Phileas, Philip Trueman, Piano non troppo, Platte (usurped), Poeloq, Poolkris, Prashanthns, Prosfilaes, Psy guy, Psyche825, Pzavon, RTC, Radiojon, Ray Van De Walker, Razorflame, Recognizance, RedWolf, Rettetast, RexNL, Rifleman 82, Rjstott, Roberta F., Robinz16, Romanm, Sally M. Peters, Samtheboy, Sandahl, Saperaud, Sbharris, SchmuckyTheCat, Schneelocke, Scott Adler, Scoutersig, Sealbubble624, Sengkang, Shafei, Shanes, Sheitan, Shoeofdeath, Shoy, Sillybilly, Silverbackmarlin, Sionus, Sivius, Sixpence, SkippyNZ, Sl, Slakr, Smokefoot, Specter01010, Squids and Chips, StaticGull, Stephenb, Stifynsemons, Stone, Suisui, T.vanschaik, TKD, Tabletop, Tagishsimon, Tarquin, Techman224, Telf, Terrx, TheProject, Thingg, Thricecube, Tide rolls, Tobias Hoevekamp, Toddst1, Trelvis, Twirligig, Tylerd22222, UberScienceNerd, Uncle G, VMS Mosaic, Vanka5, Versus22, Vishnava, Vlad4599, Vsmith, WMMartin, WTGDMan1986, WadeSimMiser, Warren oO, Warut, Watch37264, WereSpielChequers, Whispering, Whosyourjudas, Wikipediaman2000, Williamborg, Wknight94, Wpwoodard, Yngvarr, Yyy, Zetadraconis, 750 anonymous edits
Image Sources, Licenses and Contributors file:hexagonal.svg Source: http://en.wikipedia.org/w/index.php?title=File:Hexagonal.svg License: BSD Contributors: Original uploader was Danieljamesscott at en.wikipedia file:Electron shell 027 Cobalt.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_027_Cobalt.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 file:Cobalt-sample.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Cobalt-sample.jpg License: Public Domain Contributors: Benjah-bmm27 Image:Cobalt OreUSGOV.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Cobalt_OreUSGOV.jpg License: Public Domain Contributors: Bkell, Saperaud Image:2005cobalt (mined).PNG Source: http://en.wikipedia.org/w/index.php?title=File:2005cobalt_(mined).PNG License: Public Domain Contributors: User:Anwar_saadat/bubble_maps_(FAQ) Image:Cobalt - world production trend.svg Source: http://en.wikipedia.org/w/index.php?title=File:Cobalt_-_world_production_trend.svg License: Public Domain Contributors: User:Leyo Image:bristol.blue.glass.arp.750pix.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Bristol.blue.glass.arp.750pix.jpg License: Public Domain Contributors: Joanjoc, Pieter Kuiper, Saperaud, Yarl Image:Cobalamin.svg Source: http://en.wikipedia.org/w/index.php?title=File:Cobalamin.svg License: Public Domain Contributors: User:NEUROtiker
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
10
Nickel
1
Nickel cobalt ← nickel → copper ↑ Ni ↓ Pd
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
28Ni Periodic table
Appearance lustrous, metallic and silvery with a gold tinge
General Name, symbol, numberElement
Nickel
2
categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
1783
1950
2154
2410
2741
3184
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1753.0 kJ·mol−1 3rd: 3395 kJ·mol−1Atomic radiusCovalent radiusVan der Waals radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusPoisson ratioMohs hardnessVickers hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of nickel iso 56
Ni
N.A. syn
half-life 6.075 d
58
68.077%
58
59
syn
76000 y
60
26.233%
60
61
1.14%
61
62
3.634%
62
63
syn
100.1 y
64
0.926%
64
Ni Ni Ni Ni Ni Ni Ni
DM
DE (MeV)
DP
ε
-
56
γ
0.158, 0.811
-
-
59
0.0669
63
Co
Ni is stable with 30 neutron ε
Co
Ni is stable with 32 neutron Ni is stable with 33 neutron Ni is stable with 34 neutron β−
Cu
Ni is stable with 36 neutron
nickel, Ni, 28 transition metal10, 4, d58.6934(2) g·mol−1 [Ar] 4s1 3d9 2, 8, 17, 1 (Image) solid 8.908 g·cm−3 7.81 g·cm−3 1728 K,1453 °C,2651 °F 3186 K,2732 °C,5275 °F 17.48 kJ·mol−1 377.5 kJ·mol−1 (25 °C) 26.07 J·mol−1·K−1 4[1] , 3, 2, 1 [2] , -1 (mildly basic oxide) 1.91 (Pauling scale) 1st: 737.1 kJ·mol−1124 pm124±4 pm 163 pm face-centered cubic ferromagnetic (20 °C) 69.3 nΩ·m (300 K) 90.9 W·m−1·K−1 (25 °C) 13.4 µm·m−1·K−1 (r.t.) 4900 m·s−1 200 GPa 76 GPa 180 GPa 0.31 4.0 638 MPa 700 MPa 7440-02-0 Nickel (pronounced /ˈnɪkəl/) is a chemical element, with the chemical symbol Ni and atomic number 28. It is a silvery-white lustrous metal with a slight golden tinge. It is one of the four ferromagnetic elements at about room temperature. Its use has been traced as far back as 3500 BC, but it was first isolated and classified as a chemical element in 1751 by Axel Fredrik Cronstedt, who initially mistook its ore for a copper mineral. Its most important ore minerals are laterites, including limonite and garnierite, and pentlandite. Major production sites include Sudbury region in Canada, New Caledonia and Russia. The metal is corrosion-resistant, finding many uses in alloys, as a plating, in the manufacture of coins, magnets and common household utensils, as a catalyst for hydrogenation, and in a variety of other applications. Enzymes of certain life-forms contain nickel as an active center making the metal essential for them.
Nickel
Characteristics Nickel is a silvery-white metal with a slight golden tinge that takes a high polish. It is one of only four elements that are magnetic at or near room temperature. It belongs to the transition metals and is hard and ductile. It occurs most often in combination with sulfur and iron in pentlandite, with sulfur in millerite, with arsenic in the mineral nickeline, and with arsenic and sulfur in nickel galena.[3] Nickel is commonly found in iron meteorites as the alloys kamacite and taenite. Similar to the elements chromium, aluminium and titanium, nickel is a very reactive element, but is slow to react in air at normal temperatures and pressures due to the formation of a protective oxide surface. Due to its permanence in air and its slow rate of oxidation, it is used in coins, for plating metals such as iron and brass, for chemical apparatus, and in certain alloys such as German silver. Nickel is chiefly valuable for the alloys it forms, especially many superalloys, and particularly stainless steel. Nickel is also a naturally magnetostrictive material, meaning that in the presence of a magnetic field, the material undergoes a small change in length.[4] In the case of nickel, this change in length is negative (contraction of the material), which is known as negative magnetostriction and is on the order of 50 ppm. The most common oxidation state of nickel is +2 with several Ni complexes known. It is also thought that a +6 oxidation state may exist, however, this has not been demonstrated conclusively. The unit cell of nickel is a face centered cube with a lattice parameter of 0.352 nm giving a radius of the atom of 0.125 nm.[5] Because the ores of nickel are easily mistaken for ores of silver, understanding of this metal and its use dates to relatively recent times. However, the unintentional use of nickel is ancient, and can be traced back as far as 3500 BC. Bronzes from what is now Syria had contained up to 2% nickel.[6] Further, there are Chinese manuscripts suggesting that "white copper" (cupronickel, known as baitung) was used there between 1700 and 1400 BC. This Paktong white copper was exported to Britain as early as the 17th century, but the nickel content of this alloy was not discovered until 1822.[7] In medieval Germany, a red mineral was found in the Erzgebirge (Ore Mountains) which resembled copper ore. However, when miners were unable to extract any copper from it they blamed a mischievous sprite of German mythology, Nickel (similar to Old Nick) for besetting the copper. They called this ore Kupfernickel from the German Kupfer for copper.[8] [9] [10] [11] This ore is now known to be nickeline or niccolite, a nickel arsenide. In 1751, Baron Axel Fredrik Cronstedt was attempting to extract copper from kupfernickel and obtained instead a white metal that he named after the spirit which had given its name to the mineral, nickel.[12] In modern German, Kupfernickel or Kupfer-Nickel designates the alloy cupronickel. In the United States, the term "nickel" or "nick" was originally applied to the copper-nickel Indian cent coin introduced in 1859. Later, the name designated the three-cent coin introduced in 1865, and the following year the five-cent shield nickel appropriated the designation, which has remained ever since. Coins of pure nickel were first used in 1881 in Switzerland.[9] [13] After its discovery the only source for nickel was the rare Kupfernickel, but from 1824 on the nickel was obtained as byproduct of cobalt blue production. The first large scale producer of nickel was Norway, which exploited nickel rich pyrrhotite from 1848 on. The introduction of nickel in steel production in 1889 increased the demand for nickel and the nickel deposits of New Caledonia, which were discovered in 1865, provided most of the
3
Nickel
4
world's supply between 1875 and 1915. The discovery of the large deposits in the Sudbury Basin, Canada in 1883, in Norilsk-Talnakh, Russia in 1920 and in the Merensky Reef, South Africa in 1924 made large-scale production of nickel possible.[7]
Occurrence See also: Ore genesis, Category:Nickel minerals The bulk of the nickel mined comes from two types of ore deposits. The first are laterites where the principal ore minerals are nickeliferous limonite: (Fe, Ni)O(OH) and garnierite (a hydrous nickel silicate): (Ni, Mg)3Si2O5(OH). The second are magmatic sulfide deposits where the principal ore mineral is pentlandite: (Ni, Fe)9S8. In terms of supply, the Sudbury region of Ontario, Canada, produces about 30 percent of the world's supply of nickel. The Sudbury Basin deposit is theorized to have been created by a meteorite impact event early Widmanstätten pattern showing the in the geologic history of Earth. Russia contains about two forms of Nickel-Iron, Kamacite and 40% of the world's known resources at the Norilsk Taenite, in an octahedrite meteorite deposit in Siberia. The Russian mining company MMC Norilsk Nickel obtains the nickel and the associated palladium for world distribution. Other major deposits of nickel are found in New Caledonia, France, Australia, Cuba, and Indonesia. Deposits found in tropical areas typically consist of laterites which are produced by the intense weathering of ultramafic igneous rocks and the resulting secondary concentration of nickel bearing oxide and silicate minerals. Recently, a nickel deposit in western Turkey had been exploited, with this location being especially convenient for European smelters, steelmakers and factories. The one locality in the United States where nickel was commercially mined is Riddle, Oregon, where several square miles of nickel-bearing garnierite surface deposits are located. The mine closed in 1987.[14] [15] In 2005, Russia was the largest producer of nickel with about one-fifth world share closely followed by Canada, Australia and Indonesia, as reported by the British Geological Survey. Based on geophysical evidence, most of the nickel on Earth is postulated to be concentrated in the Earth's core. Kamacite and taenite are naturally occurring alloys of iron and nickel. For kamacite the alloy is usually in the proportion of 90:10 to 95:5 although impurities such as cobalt or carbon may be present, while for taenite the nickel content is between 20% and 65%. Kamacite and taenite occur in nickel-iron meteorites.
Nickel
5
Applications Nickel is used in many industrial and consumer products, including stainless steel, magnets, coinage, rechargeable batteries, electric guitar strings and special alloys. It is also used for plating and as a green tint in glass. Nickel is pre-eminently an alloy metal, and its chief use is in the nickel steels and nickel cast irons, of which there are many varieties. It is also widely used in many other alloys, such as nickel brasses and bronzes, and alloys with copper, chromium, aluminium, lead, cobalt, silver, and gold [16]
Nickel superalloy jet engine (RB199) turbine blade
The amounts of nickel used for various applications are 60% used for making nickel steels, 14% used in nickel-copper alloys and nickel silver, 9% used to make malleable nickel, nickel clad, Inconel and other superalloys, 6% used in plating, 3% use for nickel cast irons, 3% in heat and electric resistance alloys, such as Nichrome, 2% used for nickel brasses and bronzes with the remaining 3% of the nickel consumption in all other applications combined.[17] [18] In the laboratory, nickel is frequently used as a catalyst for hydrogenation, most often using Raney nickel, a finely divided form of the metal alloyed with aluminium which adsorbs hydrogen gas. Nickel is often used in coins, or occasionally as a substitute for decorative silver. The American 'nickel' five-cent coin is 75% copper and 25% nickel. The Canadian nickel minted at various periods between 1922-81 was 99.9% nickel, and was magnetic.[19] Various other nations have historically used and still use nickel in their coinage. Nickel is also used in fire assay as a collector of platinum group elements, as it is capable of full collection of all 6 elements, in addition to partial collection of gold. This is seen through the nature of nickel as a metal, as high throughput nickel mines may run PGE recovery (primarily platinum and palladium), such as Norilsk in Russia and the Sudbury Basin in Canada.
Extraction and purification Nickel is recovered through extractive metallurgy. Most sulfide ores have traditionally been processed using pyrometallurgical techniques to produce a matte for further refining. Recent advances in hydrometallurgy have resulted in recent nickel processing operations being developed using these processes. Most sulfide Nickel output in 2005 deposits have traditionally been processed by concentration through a froth flotation process followed by pyrometallurgical extraction. Nickel is extracted from its ores by conventional roasting and reduction processes which yield a metal of greater than 75% purity. Final purification of nickel oxides is performed via the Mond process, which increases the nickel concentrate to greater than 99.99% purity[20] . This process was patented by L. Mond and was used in South Wales in the 20th century. Nickel is reacted with carbon monoxide at around 50 °C to form volatile nickel carbonyl.
Nickel
6
Any impurities remain solid while the nickel carbonyl gas passes into a large chamber at high temperatures in which tens of thousands of nickel spheres, called pellets, are constantly stirred. The nickel carbonyl decomposes, depositing pure nickel onto the nickel spheres. Alternatively, the nickel carbonyl may be decomposed in a smaller chamber at 230 °C to create fine nickel powder. The resultant carbon monoxide is re-circulated through the process. The highly pure nickel produced by this process is known as carbonyl nickel. A second common form of refining involves the leaching of the metal matte followed by the electro-winning of the nickel from solution by plating it onto a cathode. In many stainless steel applications, 75% pure nickel can be used without further purification depending on the composition of the impurities. Nickel sulfide ores undergo flotation (differential flotation if Ni/Fe ratio is too low) and then are smelted. After producing the nickel matte, further processing is done via the Sherritt-Gordon process. First copper is removed by adding hydrogen sulfide, leaving a concentrate of only cobalt and nickel. Solvent extraction then efficiently separates the cobalt and nickel, with the final nickel concentration greater than 99%.
Compounds Nickel(II) sulfate is produced in large quantities by dissolving nickel metal or oxides in sulfuric acid. This compound is useful for electroplating nickel. Four halides are known to form nickel compounds,
Nickel sulfate crystals
these are nickel(II) fluoride, chloride, bromide, and iodide. Nickel(II) chloride is produced analogously by dissolving nickel residues in hydrochloric acid. Tetracarbonylnickel (Ni(CO)4), discovered by Ludwig Mond,[21] is a homoleptic complex of nickel with carbon monoxide. Having no net dipole moment, intermolecular forces are relatively weak, allowing this compound to be liquid at room temperature. Carbon monoxide reacts with nickel metal readily to give this compound; on heating, the complex decomposes back to nickel and carbon monoxide. This behavior is exploited in the Mond process for generating high-purity nickel.
Tetracoordinate nickel(II) takes both tetrahedral and square planar geometries. This is in contrast with the other group 10 elements, which tend to exist as square planar complexes. Bis(cyclooctadiene)nickel(0) is a useful intermediate in organometallic chemistry due to Tetracarbonyl nickel the easily displaced cod ligands. Nickel(III) oxide is used as the cathode in many rechargeable batteries, including nickel-cadmium, nickel-iron, nickel hydrogen, and nickel-metal hydride, and used by certain manufacturers in Li-ion batteries.[22]
Nickel
Isotopes Naturally occurring nickel is composed of 5 stable isotopes; 58Ni, 60Ni, 61Ni, 62Ni and 64Ni with 58Ni being the most abundant (68.077% natural abundance). 62Ni is the most stable known nuclide of all the existing elements, even exceeding the stability of 56Fe. 18 radioisotopes have been characterised with the most stable being 59Ni with a half-life of 76,000 years, 63Ni with a half-life of 100.1 years, and 56Ni with a half-life of 6.077 days. All of the remaining radioactive isotopes have half-lives that are less than 60 hours and the majority of these have half-lives that are less than 30 seconds. This element also has 1 meta state. Nickel-56 is produced in large quantities in type Ia supernovae and the shape of the light curve of these supernovae corresponds to the decay via beta radiation of nickel-56 to cobalt-56 and then to iron-56. Nickel-59 is a long-lived cosmogenic radionuclide with a half-life of 76,000 years. 59Ni has found many applications in isotope geology. 59Ni has been used to date the terrestrial age of meteorites and to determine abundances of extraterrestrial dust in ice and sediment. Nickel-60 is the daughter product of the extinct radionuclide 60Fe (half-life = 1.5 Myr). Because the extinct radionuclide 60Fe had such a long half-life, its persistence in materials in the solar system at high enough concentrations may have generated observable variations in the isotopic composition of 60Ni. Therefore, the abundance of 60Ni present in extraterrestrial material may provide insight into the origin of the solar system and its early history. Nickel-62 has the highest binding energy per nucleon of any isotope for any element (8.7946 Mev/nucleon). [23] Isotopes heavier than 62 Ni cannot be formed by nuclear fusion without losing energy. Nickel-48, discovered in 1999, is the most proton-rich heavy element isotope known. With 28 protons and 20 neutrons 48Ni is "double magic" (like 208Pb) and therefore unusually stable [24] . The isotopes of nickel range in atomic weight from 48 u (48Ni) to 78 u (78Ni). Nickel-78's half-life was recently measured to be 110 milliseconds and is believed to be an important isotope involved in supernova nucleosynthesis of elements heavier than iron.[25]
Biological role Nickel plays numerous roles in the biology of microorganisms and plants, though they were not recognized until the 1970s.[26] [27] In fact urease (an enzyme which assists in the hydrolysis of urea) contains nickel. The NiFe-hydrogenases contain nickel in addition to iron-sulfur clusters. Such [NiFe]-hydrogenases characteristically oxidise H2. A nickel-tetrapyrrole coenzyme, F430, is present in the methyl coenzyme M reductase which powers methanogenic archaea. One of the carbon monoxide dehydrogenase enzymes consists of an Fe-Ni-S cluster.[28] Other nickel-containing enzymes include a class of superoxide dismutase[29] and a glyoxalase.[30]
Toxicity Exposure to nickel metal and soluble compounds should not exceed 0.05 mg/cm³ in nickel equivalents per 40-hour work week. Nickel sulfide fume and dust is believed to be carcinogenic, and various other nickel compounds may be as well.[31] [32] Nickel carbonyl, [Ni(CO)4], is an extremely toxic gas. The toxicity of metal carbonyls is a function of both the toxicity of a metal as well as the carbonyl's ability to give off highly toxic carbon monoxide gas, and this one is no exception. It is explosive in air.[33]
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Nickel
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Sensitized individuals may show an allergy to nickel affecting their skin, also known as dermatitis. Sensitivity to nickel may also be present in patients with pompholyx. Nickel is an important cause of contact allergy, partly due to its use in jewellery intended for pierced ears.[34] Nickel allergies affecting pierced ears are often marked by itchy, red skin. Many earrings are now made nickel-free due to this problem. The amount of nickel which is allowed in products which come into contact with human skin is regulated by the European Union. In 2002 researchers found amounts of nickel being emitted by 1 and 2 Euro coins far in excess of those standards. This is believed to be due to a galvanic reaction.[35] It was voted Allergen of the Year in 2008 by the American Contact Dermatitis Society.[36]
Metal value The market price of nickel surged throughout 2006 and the early months of 2007; as of April 5, 2007, the metal was trading at 52,300 USD/tonne or 1.47 USD/oz.[37] The price subsequently fell dramatically from these peaks, and as of 19 January 2009 the metal was trading at 10,880 USD/tonne.[37] The US nickel coin contains 0.04 oz (1.25 g) of nickel, which at the April 2007 price was worth 6.5 cents, along with 3.75 grams of copper worth about 3 cents, making the metal value over 9 cents. Since the face value of a nickel is 5 cents, this made it an attractive target for melting by people wanting to sell the metals at a profit. However, the United States Mint, in anticipation of this practice, implemented new interim rules on December 14, 2006, subject to public comment for 30 days, which criminalize the melting and export of cents and nickels.[38] Violators can be punished with a fine of up to $10,000 and/or imprisoned for a maximum of five years. As of June 24, 2009 the melt value of a U.S. nickel is $0.0363145 which is less than the face value.[39]
See also • Category:Nickel alloys
External links • WebElements.com – Nickel
[40]
(also used as a reference)
References [1] M. Carnes et al. (2009). "A Stable Tetraalkyl Complex of Nickel(IV)". Angewandte Chemie International Edition 48: 3384. doi: 10.1002/anie.200804435 (http:/ / dx. doi. org/ 10. 1002/ anie. 200804435). [2] S. Pfirrmann et al. (2009). "A Dinuclear Nickel(I) Dinitrogen Complex and its Reduction in Single-Electron Steps". Angewandte Chemie International Edition 48: 3357. doi: 10.1002/anie.200805862 (http:/ / dx. doi. org/ 10. 1002/ anie. 200805862). [3] National Pollutant Inventory - Nickel and compounds Fact Sheet (http:/ / www. npi. gov. au/ database/ substance-info/ profiles/ 62. html) [4] UCLA - Magnetostrictive Materials Overview (http:/ / aml. seas. ucla. edu/ research/ areas/ magnetostrictive/ overview. htm) [5] Callister, William D. (2007). Materials Science and Engineering: An Introduction (7th ed.). John Wiley & Sons. ISBN 978-0-471-73696-7. [6] Rosenberg, Samuel J. (1968). Nickel and Its Alloys (http:/ / handle. dtic. mil/ 100. 2/ ADA381960). National Bureau of Standards. .
Nickel [7] McNeil, Ian (1990). "The Emergence of Nickel". An Encyclopaedia of the History of Technology. Taylor & Francis. pp. 96–100. ISBN 9780415013062. [8] Chambers Twentieth Century Dictionary, p888, W&R Chambers Ltd, 1977. [9] Baldwin, W. H. (1931). "The story of Nickel. I. How "Old Nick's" gnomes were outwitted.". Journal of Chemical Education 8: 1749. [10] Baldwin, W. H. (1931). "The story of Nickel. II. Nickel comes of age.". Journal of Chemical Education 8: 1954. [11] Baldwin, W. H. (1931). "The story of Nickel. III. Ore, matte, and metal.". Journal of Chemical Education 8: 2325. [12] Weeks, Mary Elvira (1932). "The discovery of the elements: III. Some eighteenth-century metals". Journal of Chemical Education 9: 22. [13] Molloy, Bill (2001-11-08). " Trends of Nickel in Coins - Past, Present and Future (http:/ / www. nidi. org/ index. cfm/ ci_id/ 160. htm)". The Nickel Institute. . Retrieved 2008-11-19. [14] " The Nickel Mountain Project (http:/ / www. oregongeology. com/ sub/ publications/ OG/ OBv15n10. pdf)". Ore Bin 15 (10): 59–66. 1953. . [15] " Environment Writer: Nickel (http:/ / www. environmentwriter. org/ resources/ backissues/ chemicals/ nickel. htm)". National Safety Council. 2006. . Retrieved 2009-01-10. [16] Davis, Joseph R. (2000). " Uses of Nickel (http:/ / books. google. de/ books?id=IePhmnbmRWkC)". ASM Specialty Handbook: Nickel, Cobalt, and Their Alloys. ASM International. pp. 7–13. ISBN 9780871706850. . [17] Kuck, Peter H.. " Mineral Commodity Summaries 2006: Nickel (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ nickel/ mcs-2008-nicke. pdf)". United States Geological Survey. . Retrieved 2008-11-19. [18] Kuck, Peter H.. " Mineral Yearbook 2006: Nickel (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ nickel/ myb1-2006-nicke. pdf)". United States Geological Survey. . Retrieved 2008-11-19. [19] " Industrious, enduring–the 5-cent coin (http:/ / www. mint. ca/ store/ mint/ learn/ circulation-currency-1100028)". Royal Canadian Mint. 2008. . Retrieved 2009-01-10. [20] Mond L, Langer K, Quincke F (1890). "Action of carbon monoxide on nickel". Journal of the Chemical Society 57: 749–753. doi: 10.1039/CT8905700749 (http:/ / dx. doi. org/ 10. 1039/ CT8905700749). [21] Mond L, Langer K, Quincke F (1890). "Action of carbon monoxide on nickel". Journal of the Chemical Society 57: 749–753. doi: 10.1039/CT8905700749 (http:/ / dx. doi. org/ 10. 1039/ CT8905700749). [22] http:/ / www. greencarcongress. com/ 2008/ 12/ imara-corporati. html [23] " The Most Tightly Bound Nuclei (http:/ / hyperphysics. phy-astr. gsu. edu/ hbase/ nucene/ nucbin2. html#c1)". . Retrieved 2008-11-19. [24] W., P. (October 23, 1999). " Twice-magic metal makes its debut - isotope of nickel (http:/ / www. findarticles. com/ p/ articles/ mi_m1200/ is_17_156/ ai_57799535)". Science News. . Retrieved 2006-09-29. [25] Castelvecchi, Davide (2005-04-22). " Atom Smashers Shed Light on Supernovae, Big Bang (http:/ / skyandtelescope. com/ news/ article_1502_1. asp)". . Retrieved 2008-11-19. [26] Astrid Sigel, Helmut Sigel and Roland K.O. Sigel, ed (2008). Nickel and Its Surprising Impact in Nature. Metal Ions in Life Sciences. 2. Wiley. ISBN 978-0-470-01671-8. [27] Hausinger, R. P. (1987). " Nickel utilization by microorganisms (http:/ / www. pubmedcentral. nih. gov/ picrender. fcgi?artid=373090& blobtype=pdf)". Microbiol Review 51 (1): 22–42. . [28] Jaouen, G. (2006). Bioorganometallics: Biomolecules, Labeling, Medicine. Wiley-VCH: Weinheim. [29] Szilagyi, R. K.; Bryngelson, P. A.; Maroney, M. J.; Hedman, B.; Hodgson, K. O.; Solomon, E. I. (2004). "S K-Edge X-ray Absorption Spectroscopic Investigation of the Ni-Containing Superoxide Dismutase Active Site: New Structural Insight into the Mechanism". Journal of the American Chemical Society 126 (10): 3018–3019. doi: 10.1021/ja039106v (http:/ / dx. doi. org/ 10. 1021/ ja039106v). [30] Thornalley, P. J. (2003). " Glyoxalase I--structure, function and a critical role in the enzymatic defence against glycation (http:/ / www. biochemsoctrans. org/ bst/ 031/ bst0311343. htm)". Biochemical Society Transactions 31: 1343–1348. doi: 10.1042/BST0311343 (http:/ / dx. doi. org/ 10. 1042/ BST0311343). . [31] KS Kasprzak, FW Sunderman Jr, K Salnikow. Nickel carcinogenesis. Mutation Research. 2003 December 10;533(1-2):67-97. PubMed (http:/ / www. ncbi. nlm. nih. gov/ sites/ entrez?cmd=retrieve& db=pubmed& list_uids=14643413& dopt=Abstract) [32] JK Dunnick, MR Elwell, AE Radovsky, JM Benson, FF Hahn, KJ Nikula, EB Barr, CH Hobbs. Comparative Carcinogenic Effects of Nickel Subsulfide, Nickel Oxide, or Nickel Sulfate Hexahydrate Chronic Exposures in the Lung. Cancer Research. 1995 November 15;55(22):5251-6. PubMed (http:/ / www. ncbi. nlm. nih. gov/ sites/ entrez?cmd=retrieve& db=pubmed& list_uids=7585584& dopt=Abstract) [33] http:/ / msds. chem. ox. ac. uk/ NI/ nickel_carbonyl. html [34] Thyssen JP, Linneberg A, Menné T, Johansen JD (2007). " The epidemiology of contact allergy in the general population—prevalence and main findings (http:/ / www. blackwell-synergy. com/ doi/ full/ 10. 1111/ j. 1600-0536. 2007. 01220. x)". Contact Dermatitis 57 (5): 287–99. doi: 10.1111/j.1600-0536.2007.01220.x (http:/ / dx. doi. org/ 10. 1111/ j. 1600-0536. 2007. 01220. x). PMID 17937743. .
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Nickel [35] Nestle, O.; Speidel, H.; Speidel, M. O. (2002). " free abstract High nickel release from 1- and 2-euro coins (http:/ / www. nature. com/ nature/ journal/ v419/ n6903/ abs/ 419132a. html)". Nature 419: 132. doi: 10.1038/419132a (http:/ / dx. doi. org/ 10. 1038/ 419132a). free abstract. [36] " Nickel Named 2008 Contact Allergen of the Year (http:/ / www. nickelallergyinformation. com/ 2008/ 06/ nickel-named-2008-contact-alle. htm)". . Retrieved 2009-06-06. [37] " LME nickel price graphs (http:/ / www. lme. co. uk/ nickel_graphs. asp)". London Metal Exchange. . Retrieved 2009-06-06. [38] United States Mint Moves to Limit Exportation & Melting of Coins (http:/ / www. usmint. gov/ pressroom/ index. cfm?action=press_release& ID=724), The United States Mint, press release, December 14, 2006 [39] " United States Circulating Coinage Intrinsic Value Table (http:/ / www. coinflation. com/ )". Coininflation.com. . Retrieved 2009-06-06. [40] http:/ / www. webelements. com/ webelements/ elements/ text/ Ni/ index. html
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Article Sources and Contributors
Article Sources and Contributors Nickel Source: http://en.wikipedia.org/w/index.php?oldid=306772057 Contributors: 293.xx.xxx.xx, 56, A D 13, Abce2, Academic Challenger, Adam Johnston, Adashiel, Addshore, Ahoerstemeier, Aitias, Alansohn, Ale jrb, Algebra, Alias Flood, Alpha 4615, AndonicO, Andres, Andy M. Wang, Andycjp, Animum, Anomalocaris, Anonymous editor, Antandrus, Antonio Lopez, Anwar saadat, Arakunem, Archimerged, Arthurzahn, Ascidian, Aunt-jemima, Aussie Alchemist, B.d.mills, BANZ111, Babajobu, Bahuston, Bambookoloa, BarretBonden, Basicdesign, Bcorr, Bearcat, Beetstra, BenFrantzDale, Benbest, BjKa, Bkell, Blow of Light, BlueEarth, Bluh10, Bobo192, Bovineone, Brettford23, Brian Huffman, Brian0918, BrokenSegue, Bryan Derksen, Bsmithme, C.A.T.S. CEO, CWii, Cadmium, Cadwallader, Caltas, Can't sleep, clown will eat me, CanadianLinuxUser, CardinalDan, Carnildo, Carrotsrgangsta, Catalyst2007, Catgut, Causa sui, Caz1234567, CesarB, Chairman S., Chameleon, ChemGardener, ChemNerd, Choess, Chris 73, ChrisGriswold, ChrisO, Coemgenus, Cometstyles, Conversion script, CopperKettle, Coppertwig, CorvetteZ51, Curps, CyanideSandwich, Cybernetic, Cynical, D. Recorder, DBragagnolo, DMacks, Dan Guan, DanielCD, Dark Mage, DarkFalls, Darrendeng, Darrien, Daven200520, David Latapie, Davy p, Dawn Bard, Daz0123, Deadlock, Dekisugi, Deli nk, Delicious carbuncle, Delta G, DennyColt, DerHexer, Discospinster, Dissident, Dmohs, Doonhamer, Download, Drini, Dspradau, Durin, Dwane E Anderson, Dwmyers, Dycedarg, Edgar181, El C, Eldin raigmore, Electrostatic1, Elliotthedunk, Ellisosee, Emperorbma, Epbr123, Eric-Wester, Eubulides, Evcana, Everyking, Falconsbeast07, Femto, Flagmantho, Flaresam, Flewis, Flypro, Fredbauder, Fredrik, Fritzpoll, Gadfium, Gail, Gene Nygaard, Giftlite, Gilliam, Gman124, Gogo Dodo, Goldman02, Golthar, Greenman, Grendelkhan, Groyolo, Gzkn, Haakon, Hadal, Haham hanuka, Hairy Dude, Hda3ku, Heirpixel, Hellsuxx1, Hephaestos, Hexagon1, Hike395, Husond, Hut 8.5, I3lizzard, IanOsgood, Icairns, Icek, Iepeulas, Ihope4me, Ikalwebshow, Inferno, Lord of Penguins, Ixfd64, J.delanoy, JForget, JaGa, Jacks216, Janke, Jaraalbe, Jaxl, Jeddtheonedick, Jennavecia, JeremyA, Jerzy, Jim Swenson, Jimothytrotter, Jimp, Jo Stainless, Joanjoc, John, John254, Johnleemk, JohnnyRush10, Jons63, Jorge Stolfi, Jose77, Journeyman, Jrockley, Jrvz, Jtkiefer, Juanamac, Junglecat, Kaitlyn18, Karl-Henner, Karlhahn, Keilana, Kelisi, Kevin Breitenstein, Kevink0001, Khelben "Blackstaff" Arunsun, Kimshady, King of Hearts, Kingpin13, Knowledge Seeker, KnowledgeOfSelf, Knutux, Kungfuadam, Kuru, Kwamikagami, LA2, Lcarscad, Lewismeurig, Lightmouse, Ligulem, Looxix, Lucasbfr, LuigiManiac, Luna Santin, MER-C, MZMcBride, Madhero88, MadisonMichael, Magnus Manske, Malerin, Mangojuice, MarkSutton, Markaci, Master of Puppets, Materialscientist, Matt Beard, Matthead, Mav, McTrixie, Melloss, Mgimpel, Michael Hardy, Mike Dillon, Mikeblas, Minesweeper, Mion, Miquonranger03, Miss Madeline, Mmoneypenny, Monkeynoze, Mosterman88, Mr Stephen, Mr.Z-man, Mrstone56, Mschel, Nae'blis, Nanosilver, NawlinWiki, Ncmvocalist, Neutrality, NewEnglandYankee, Nibuod, Nick Y., Nihiltres, Nilfanion, Nuttycoconut, Oda Mari, OldakQuill, Ossmann, Overand, Owen, Oxymoron83, PCock, Pagingmrherman, Pakaran, Paul Erik, PaulHanson, Payo, Peter Deer, Peterckw, Petergans, Peterlewis, Pgk, Philip Trueman, Physchim62, Piano non troppo, Pizza1512, Poolkris, Possum, Pras, Prashanthns, Prophet121, Psyche825, Purgatory Fubar, Queerr00, Quinsareth, Quintote, Qweedsa, RTC, Ra'ike, Radiojon, RainR, RaseaC, Ray Chason, RayAYang, Red King, Remember, Res2216firestar, Reyk, Rich Farmbrough, Rifleman 82, Rjanag, Rjwilmsi, Roadrunner, RobLa, Roberta F., Rolinator, Romanm, Roni2204, Rossnorman, Rursus, S234432, SD6-Agent, Sam Hocevar, Sam Korn, Sanders muc, Sannse, Saperaud, Sbharris, Scaife, Scetoaux, Schneelocke, Sengkang, Seth Ilys, Shafei, Shanes, Sheitan, Shoy, Sikh123123, Silsor, Silverbackmarlin, SimonP, Sirex98, Sjakkalle, Skeptical chymist, Skwa, Sl, Slakr, Slysplace, Smokefoot, Snowolf, Spinningspark, Splash, Spp143, Stanleypark, Sterence96, Stismail, Stone, StuffOfInterest, Suisui, Sunborn, Superfreaky56, Switchercat, Swpb, Tagishsimon, The Last Heretic, The Obento Musubi, The way, the truth, and the light, TheSuave, Thehelpfulone, Theseeker4, Thingg, Thricecube, Tiddly Tom, Tide rolls, Toyoumygirl, Tucker111111, Turgan, UberScienceNerd, UnitedStatesian, Unschool, Until It Sleeps, Urhixidur, Userafw, Username314, Vanka5, Vashtihorvat, Velela, Vivio Testarossa, Voyager3, Vsmith, Waliko, Walkerma, Where, Why Not A Duck, Wik, William Ortiz, WilliamH, Wimt, Winchelsea, Windchaser, Yath, Yeonhum1357, Yonatan, Yyy, Zen611, Zenlax, Zoicon5, 1003 anonymous edits
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11
Copper
1
Copper nickel ← copper → zinc ↑ Cu ↓ Ag
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HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
29Cu Periodic table
Appearance reddish/orange metallic lustre
metallic copper General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
1509
1661
1850
2089
2404
2834
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1957.9 kJ·mol−1 3rd: 3555 kJ·mol−1Atomic radiusCovalent radiusVan der Waals
Copper
2
radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusPoisson ratioMohs hardnessVickers hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of copper iso
N.A.
half-life
63
69.15%
63
65
30.85%
65
Cu Cu
DM
DE (MeV)
DP
Cu is stable with 34 neutron Cu is stable with 36 neutron
copper, Cu, 29 transition metal11, 4, d63.546(3) g·mol−1 [Ar] 3d10 4s1 2, 8, 18, 1 (Image) solid 8.94 g·cm−3 8.02 g·cm−3 1357.77 K,1084.62 °C,1984.32 °F 2835 K,2562 °C,4643 °F 13.26 kJ·mol−1 300.4 kJ·mol−1 (25 °C) 24.440 J·mol−1·K−1 +1, +2, +3, +4 (mildly basic oxide) 1.90 (Pauling scale) 1st: 745.5 kJ·mol−1128 pm132±4 pm 140 pm face-centered cubic diamagnetic (20 °C) 16.78 nΩ·m (300 K) 401 W·m−1·K−1 (25 °C) 16.5 µm·m−1·K−1 (r.t.) (annealed) 3810 m·s−1 110–128 GPa 48 GPa 140 GPa 0.34 3.0 369 MPa 874 MPa 7440-50-8 Copper (pronounced /ˈkɒpər/) is a chemical element with the symbol Cu (Latin: cuprum) and atomic number 29. It is a ductile metal with very high thermal and electrical conductivity. Pure copper is rather soft and malleable and a freshly-exposed surface has a pinkish or peachy color. It is used as a thermal conductor, an electrical conductor, a building material, and a constituent of various metal alloys. Copper metal and alloys have been used for thousands of years. In the Roman era, copper was principally mined on Cyprus, hence the origin of the name of the copper disc made by continuous metal as Cyprium, "metal of Cyprus", later shortened to casting, etched Cuprum. There may be insufficient reserves to sustain current high rates of copper consumption.[1] Some countries, such as Chile and the United States, still have sizable reserves of the metal which are extracted through large open pit mines. Copper compounds are known in several oxidation states, usually 2+, where they often impart blue or green colors to natural minerals such as turquoise and have been used historically widely as pigments. Copper as both metal and pigmented salt, has a significant presence in decorative art. Copper 2+ ions are soluble in water, where they function at low concentration as bacteriostatic substances and fungicides. For this reason, copper metal can be used as an anti-germ surface that can add to the anti-bacterial and antimicrobial features of buildings such as hospitals.[2] In sufficient amounts, copper salts can be poisonous to higher organisms as well. However, despite universal toxicity at high concentrations, the 2+ copper ion at lower concentrations is an essential trace nutrient to all higher plant and animal life. In animals, including humans, it is found widely in tissues, with concentration in liver, muscle, and bone. It functions as a co-factor in various enzymes and in copper-based pigments.
Copper
3
History Copper Age Copper, as native copper, is one of the few metals to occur naturally as an un-compounded mineral. Copper was known to some of the oldest civilizations on record, and has a history of use that is at least 10,000 years old. Some estimates of copper's discovery place this event around 9000 BC in the Middle East.[3] A copper pendant was found in what is now northern Iraq that dates to 8700 BC.[4] It is probable that gold and meteoritic iron were the only metals used by humans before copper.[5] By 5000 BC, there are signs of copper smelting: the refining of copper from simple copper compounds such as malachite or azurite. Among archaeological sites in Anatolia, Çatal Höyük (~6000 BC) features native copper artifacts and smelted lead beads, but no smelted copper. Can Hasan (~5000 BC) had access to smelted copper but the oldest smelted copper artifact found (a copper chisel from the chalcolithic site of Prokuplje in Serbia) has pre-dated Can Hasan by 500 years. The smelting facilities in the Balkans appear to be more advanced than the Turkish forges found at a later date, so it is quite probable that copper smelting originated in the Balkans. Investment casting was realized in 4500-4000 BCE in Southeast Asia.[3] Copper smelting appears to have been developed independently in several parts of the world. In addition to its development in the Balkans by 5500 BC, it was developed in China before 2800 BC, in the Andes around 2000 BC, in Central America around 600 AD, and in West Africa around 900 AD.[6] Copper is found extensively in the Indus Valley Civilization by the 3rd millennium BC. In Europe, Ötzi the Iceman, a well-preserved male Ancient Copper ingot from Zakros, Crete is dated to 3300-3200 BC, was found with an axe with shaped in the form of an animal skin typical a copper head 99.7% pure. High levels of arsenic in for that era. his hair suggest he was involved in copper smelting. Over the course of centuries, experience with copper has assisted the development of other metals; for example, knowledge of copper smelting led to the discovery of iron smelting. In the Americas production in the Old Copper Complex, located in present day Michigan and Wisconsin, was dated back to between 6000 to 3000 BC.[7]
Bronze Age Alloying of copper with zinc or tin to make brass or bronze was practiced soon after the discovery of copper itself. There exist copper and bronze artifacts from Sumerian cities that date to 3000 BC,[8] and Egyptian artifacts of copper and copper-tin alloys nearly as old. In one pyramid, a copper plumbing system was found that is 5000 years old.[9] The Egyptians found that adding a small amount of tin made the metal easier to cast, so copper-tin (bronze) alloys were found in Egypt almost as soon as copper was found. Very important sources of copper in the Levant were located in Timna valley (Negev, now in southern Israel) and Faynan (biblical Punon, Jordan).[10] By 2000 BC, Europe was using bronze.[8] The use of bronze became so widespread in Europe approximately from 2500 BC to 600 BC that it has been named the Bronze Age. The transitional period in certain regions between the preceding Neolithic period and the
Copper
4
Bronze Age is termed the Chalcolithic ("copper-stone"), with some high-purity copper tools being used alongside stone tools. Brass (copper-zinc alloy) was known to the Greeks, but only became a significant supplement to bronze during the Roman empire. During the Bronze Age, one copper mine at Great Orme in North Wales, extended for a depth of 70 meters.[11] At Alderley Edge in Cheshire, carbon dates have established mining at around 2280 to 1890 BC (at 95% probability).[12]
Antiquity and Middle Ages
In alchemy the symbol for copper, perhaps a stylized mirror, was also the symbol for the goddess and planet Venus.
In Greek the metal was known by the name chalkos (χαλκός). Copper was a very important resource for the Romans, Greeks and other ancient peoples. In Roman times, it became known as aes Cyprium (aes being the generic Latin term for copper alloys such as bronze and other metals, and Cyprium because so much of it was mined in Cyprus). From this, the phrase was simplified to cuprum, hence the English copper. Copper was associated with the goddess Aphrodite/Venus in mythology and alchemy, owing to its lustrous beauty, its ancient use in producing mirrors, and its association with Cyprus, which was sacred to the goddess. In astrology, alchemy the seven heavenly bodies known to the ancients were associated with seven metals also known in antiquity, and Venus was assigned to copper.[13] Britain's first use of brass occurred around the 3rd -
Chalcolithic copper mine in Timna Valley, Negev Desert, Israel.
2nd century B.C. In North America, copper mining began with marginal workings by Native Americans. Native copper is known to have been extracted from sites on Isle Royale with primitive stone tools between 800 and 1600.[14]
Copper metallurgy was flourishing in South America, particularly in Peru around the beginning of the first millennium AD. Copper technology proceeded at a much slower rate on other continents. Africa's major location for copper reserves is Zambia. Copper burial ornamentals dated from the 15th century have been uncovered, but the metal's commercial production did not start until the early 1900s. Australian copper artifacts exist, but they appear only after the arrival of the Europeans; the aboriginal culture apparently did not develop their own metallurgical abilities. Crucial in the metallurgical and technological worlds, copper has also played an important cultural role, particularly in currency. Romans in the 6th through 3rd centuries B.C. used copper lumps as money. At first, just the copper itself was valued, but gradually the shape and look of the copper became more important. Julius Caesar had his own coins, made from a copper-zinc alloy, while Octavianus Augustus Caesar's) coins were made from Cu-Pb-Sn alloys.
Copper
5
The gates of the Temple of Jerusalem used Corinthian bronze made by depletion gilding. Corinthian bronze was most prevalent in Alexandria, where alchemy is thought to have begun.[15] In ancient India (before 1000 B.C.), copper was used in the holistic medical science Ayurveda for surgical instruments and other medical equipment. Ancient Egyptians (~2400 B.C.) used copper for sterilizing wounds and drinking water, and as time passed, (~1500 B.C.) for headaches, burns, and itching. Hippocrates (~400 B.C.) used copper to treat leg ulcers associated with varicose veins. Ancient Aztecs fought sore throats by gargling with copper mixtures. Copper is also the part of many rich stories and legends, such as that of Iraq's Baghdad Battery. Copper cylinders soldered to lead, which date back to 248 B.C. to 226 A.D, resemble a galvanic cell, leading people to believe this may have been the first battery. This claim has so far not been substantiated. The Bible also refers to the importance of copper: "Men know how to mine silver and refine gold, to dig iron from the earth and melt copper from stone" (Job. 28:1-2).
Modern period Throughout history, copper's use in art has extended far beyond currency. Vannoccio Biringuccio, Giorgio Vasari and Benvenuto Cellini are three Renaissance sculptors from the mid 1500s, notable for their work with bronze. From about 1560 to about 1775, thin sheets of copper were commonly used as a canvas for paintings. Silver plated copper was used in the pre-photograph known as the daguerreotype. The Statue of Liberty, dedicated on October 28, 1886, was constructed of copper thought to have come from French-owned mines in Norway.
Miners at the Tamarack Mine in Copper Country, Michigan, USA in 1905
Plating was a technology that began started in the mid 1600s in some areas. One common use for copper plating, widespread in the 1700s, was the sheathing of ships' hulls. Copper sheathing could be used to protect wooden hulled ships from algae, and from the shipworm "toredo", a saltwater clam. The ships of Christopher Columbus were among the earliest to have this protection.[16] In 1801 Paul Revere established America's first copper rolling mill in Canton, Massachusetts. In the early 1800s, it was discovered that copper wire could be used as a conductor, but it wasn't until 1990 that copper, in oxide form, was discovered for use as a superconducting material. The German scientist Gottfried Osann invented powder metallurgy of copper in 1830 while determining the metal's atomic weight. Around then it was also discovered that the amount and type of alloying element (e.g. tin) would affect the tones of bells, allowing for a variety of rich sounds, leading to bell casting, another common use for copper and its alloys. The Norddeutsche Affinerie in Hamburg was the first modern electroplating plant starting its production in 1876.[17] Flash smelting, was developed by Outokumpu in Finland and first applied at the Harjavalta plant in 1949. The process makes smelting more energy efficient and is today used for 50% of the world’s primary copper production.[18]
Copper Copper has been pivotal in the economic and sociological worlds, notably disputes involving copper mines. The 1906 Cananea Strike in Mexico dealt with issues of work organization. The Teniente copper mine (1904-1951) raised political issues about capitalism and class structure. Japan's largest copper mine, the Ashio mine, was the site of a riot in 1907. The Arizona miners' strike of 1938 dealt with American labor issues including the "right to strike".
Characteristics Color Copper has a reddish, orangish, or brownish color because a thin layer of tarnish (including oxides) gradually forms on its surface when gases (especially oxygen) in the air react with it. But pure copper, when fresh, is actually a pinkish or peachy metal. Copper, caesium and gold are the only three elemental metals with a natural color other than gray or silver.[19] The usual gray color of metals depends on their "electron sea" that is capable of absorbing and re-emitting Copper just above its melting point keeps its pink luster color photons over a wide range of frequencies. Copper has its when enough light outshines the characteristic color because of its unique band structure. By orange incandescence color. Madelung's rule the 4s subshell should be filled before electrons are placed in the 3d subshell but copper is an exception to the rule with only one electron in the 4s subshell instead of two. The energy of a photon of blue or violet light is sufficient for a d band electron to absorb it and transition to the half-full s band. Thus the light reflected by copper is missing some blue/violet components and appears red. This phenomenon is shared with gold which has a corresponding 5s/4d structure.[20] In its liquefied state, a pure copper surface without ambient light appears somewhat greenish, a characteristic shared with gold. When liquid copper is in bright ambient light, it retains some of its pinkish luster. When copper is burnt in oxygen it gives off a black oxide.
Group 11 of the periodic table Copper occupies the same family of the periodic table as silver and gold, since they each have one s-orbital electron on top of a filled electron shell which forms metallic bonds. This similarity in electron structure makes them similar in many characteristics. All have very high thermal and electrical conductivity, and all are malleable metals. Among pure metals at room temperature, copper has the second highest electrical and thermal conductivity, after silver.[21]
6
Copper
7
Occurrence Copper can be found as native copper in mineral form (for example, in Michigan's Keewenaw Peninsula). It is a polycrystal, with the largest single crystals measuring 4.4x3.2x3.2 cm3.[22] Minerals such as the sulfides: chalcopyrite (CuFeS2), bornite (Cu5FeS4), covellite (CuS), chalcocite (Cu2S) are sources of copper, as are the carbonates: azurite (Cu3(CO3)2(OH)2) and malachite (Cu2CO3(OH)2) and the oxide: cuprite (Cu2O).[21]
Mechanical properties
Native copper (~ 4 cm in size)
Copper is easily worked, being both ductile and malleable. The ease with which it can be drawn into wire makes it useful for electrical work in addition to its excellent electrical properties. Copper can be machined, although it is usually necessary to use an alloy for intricate parts, such as threaded components, to get really good machinability characteristics. Good thermal conduction makes it useful for heatsinks and in heat exchangers. Copper has good corrosion resistance, but not as good as gold. It has excellent brazing and soldering properties and can also be welded, although best results are obtained with gas metal arc welding.[23] Copper is normally supplied, as with nearly all metals for industrial and commercial use, in a fine grained polycrystalline form. Polycrystalline metals have greater strength than monocrystalline forms, and the difference is greater for smaller grain (crystal) sizes. The reason is due to the inability of stress dislocations in the crystal structure to cross the grain boundaries.[24]
Electrical properties At 59.6 × 106 S/m copper has the second highest electrical conductivity of any element, just after silver. This high value is due to virtually all the valence electrons (one per atom) taking part in conduction. The resulting free electrons in the copper amount to a huge charge density of 13.6x109 C/m3. This high charge density is responsible for the rather slow drift velocity of currents in copper cable (drift velocity may be calculated as the ratio of current density to charge density). For instance, at a current density of 5x106 A/m2 (typically, the maximum current density present in household wiring and grid distribution) the drift velocity is just a little over ⅓ mm/s.[25]
Corrosion
Copper electrical busbars distributing power to a large building.
Copper
8
Pure water and air/oxygen Copper is a metal that does not react with water (H2O), but the oxygen of the air will react slowly at room temperature to form a layer of brown-black copper oxide on copper metal.
The Pourbaix diagram for copper in pure water, or acidic or alkali conditions. It can be seen that copper in neutral water is more noble than hydrogen.
It is important to note that in contrast to the oxidation of iron by wet air that the layer formed by the reaction of air with copper has a protective effect against further corrosion. On old copper roofs a green layer of copper carbonate, called verdigris, can often be seen. A notable example of this is on the Statue of Liberty. In contact with other metals
Copper should not be in direct mechanical contact with metals of different electropotential (for example, a copper pipe joined to an iron pipe), especially in the presence of moisture, as the completion of an electrical circuit (for instance through the common ground) will cause the juncture to act as an electrochemical cell (like a single cell of a battery). The weak electrical currents themselves are harmless but the electrochemical reaction will cause the conversion of the iron to other compounds, eventually destroying the functionality of the union. This problem is usually solved in plumbing by separating copper pipe from iron pipe with some non-conducting segment (usually plastic or rubber). Sulfide media Copper
metal
reacts
with
hydrogen
sulfide-containing solutions, forming different copper sulfides on its surface.
sulfidea
series
and of
The Pourbaix diagram is very complex due to the existence of many different sulfides. In sulfide-containing solutions copper is less noble than hydrogen and will corrode. This can be observed in everyday life when copper metal surfaces tarnish after exposure to air containing sulfur compounds.
The Pourbaix diagram for copper in water containing sulfide
Copper
9
Ammonia media Copper is slowly dissolved in oxygen-containing ammonia solutions because the ammonia forms water-soluble copper complexes. The formation of these complexes causes the corrosion to become more thermodynamically favored than the corrosion of copper in an identical solution that does not contain the ammonia.
The Pourbaix diagram for copper in 10 M ammonia solution
Chloride media Copper reacts with a combination of oxygen and hydrochloric acid to form a series of copper chlorides. Copper(II) chloride (green/blue) when boiled with copper metal undergoes a symproportionation reaction to form white copper(I) chloride.
Germicidal effect Copper is germicidal, via the oligodynamic effect. For example, brass doorknobs disinfect themselves of many bacteria within a period of eight hours.[26] Antimicrobial properties of copper are effective against MRSA,[27] Escherichia coli[28] and other pathogens.[29] longer time is required to kill bacteria.
The Pourbaix diagram for copper in a chloride solution [30] [31]
In colder temperature,
Copper has the intrinsic ability to kill a variety of potentially harmful pathogens. On February 29, 2008, the United States EPA registered 275 alloys, containing greater than 65% nominal copper content, as antimicrobial materials[32] . Registered alloys include pure copper, an assortment of brasses and bronzes, and additional alloys. EPA-sanctioned tests using Good Laboratory Practices were conducted in order to obtain several antimicrobial claims valid against: methicillin-resistant Staphylococcus aureus (MRSA), Enterobacter aerogenes, Escherichia coli O157: H7 and Pseudomonas aeruginosa. The EPA registration allows the manufacturers of these copper alloys to legally make public health claims as to the health effects of these materials. Several of the aforementioned bacteria are responsible for a large portion of the nearly two million hospital-acquired infections contracted each year in the United States[33] . Frequently touched surfaces in hospitals and public facilities harbor bacteria and increase the risk for contracting infections. Covering touch surfaces with copper alloys can help reduce microbial contamination associated with hospital-acquired-infections on these surfaces.
Copper
10
Isotopes Copper has 29 distinct isotopes ranging in atomic mass from 52 to 80. Two of these, 63Cu and 65Cu, are stable and occur naturally, with 63Cu comprising approximately 69% of naturally occurring copper.[34] The other 27 isotopes are radioactive and do not occur naturally. The most stable of these is 67 Cu with a half-life of 61.83 hours. The least stable is 54Cu with a half-life of approximately 75 ns. Unstable copper isotopes with atomic masses below 63 tend to undergo β+ decay, while isotopes with atomic masses above 65 tend to undergo β− decay. 64Cu decays by both β+ and β−.[34] 68
Cu, 69Cu, 71Cu, 72Cu, and 76Cu each have one metastable isomer. 70Cu has two isomers, making a total of 7 distinct isomers. The most stable of these is 68mCu with a half-life of 3.75 minutes. The least stable is 69mCu with a half-life of 360 ns.[34]
Production Output Most copper ore is mined or extracted as copper
Chuquicamata (Chile). The largest open pit copper mines in the world.
sulfides from large open pit mines in porphyry copper deposits that contain 0.4 to 1.0 percent copper. Examples include: Chuquicamata in Chile and El Chino Mine in New Mexico. The average abundance of copper found within crustal rocks is approximately 68 ppm by mass, and 22 ppm by atoms. In 2005, Chile was the top mine producer of copper with at least one-third world share followed by the USA, Indonesia and Peru, reports the British Geological Survey.[21]
Copper output in 2005
Reserves Copper has been in use at least 10,000 years, but more than 95 percent of all copper ever mined and smelted has been extracted since 1900. Like fossil fuels, copper is a finite resource. The Earth has an estimated 61 years of copper reserves remaining. Environmental analyst, Lester Brown, however, has suggested copper might run out within 25 years based on a reasonable extrapolation of 2% growth per year.[35] World production trend
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As consumption in India and China increases, copper supplies are becoming scarcer.[36] The copper price has quintupled from the 60-year low in 1999, rising from US$0.60 per pound (US$1.32/kg) in June 1999 to US$3.75 per pound (US$8.27/kg) in May 2006, where it dropped to US$2.40 per pound (US$5.29/kg) in February 2007 then rebounded to US$3.50 per pound (US$7.71/kg = £3.89 = €5.00) in April 2007.[37] By Copper Prices 2003 - 2008 in USD early February 2009, however, weakening global demand and a steep fall in commodity prices since the previous year's highs had left copper prices at US$1.51 per pound.[38] Production is expected to reach a peak at some point in the future and decline thereafter, according to the Hubbert peak theory. The Intergovernmental Council of Copper Exporting Countries (CIPEC), defunct since 1992, once tried to play a similar role for copper as OPEC does for oil, but never achieved the same influence, not least because the second-largest producer, the United States, was never a member. Formed in 1967, its principal members were Chile, Peru, Zaire, and Zambia.
Applications Copper is malleable and ductile and is a good conductor of both heat and electricity. The purity of copper is expressed as 4N for 99.99% pure or 7N for 99.99999% pure. The numeral gives the number of nines after the decimal point when expressed as a decimal (e.g. 4N means 0.9999, or 99.99%). Copper is often too soft for its applications, so it is incorporated in numerous alloys. For example, brass is a copper-zinc alloy, and bronze is a copper-tin alloy.[39] It is used extensively, in products such as:
Piping • including water supply. • used extensively in refrigeration and air conditioning equipment because of its ease of fabrication and soldering, as well as high conductivity to heat.
Electrical applications • Copper wire. • Oxygen-free copper. • Electromagnets.
Assorted copper fittings.
• Printed circuit boards. • Lead free solder, alloyed with tin. • Electrical machines, especially electromagnetic motors, generators and transformers. • Electrical relays, electrical busbars and electrical switches. • Vacuum tubes, cathode ray tubes, and the magnetrons in microwave ovens. • Wave guides for microwave radiation.
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• Integrated circuits, increasingly replacing aluminium because of its superior electrical conductivity. • As a material in the manufacture of computer heat sinks, as a result of its superior heat dissipation capacity to aluminium.
Architecture / Industry • Copper has been used as water-proof roofing material since ancient times, giving many old buildings their greenish roofs and domes. Initially copper oxide forms, replaced by cuprous and cupric sulfide, and finally by copper carbonate. The final carbonate patina (termed verdigris) is highly resistant to corrosion.[40] • Statuary: The Statue of Liberty, for example, contains 179,220 pounds (81.3 tonnes) of copper. • Alloyed with nickel, e.g. cupronickel and Monel, used as corrosive resistant materials in shipbuilding.
Copper roof on the Minneapolis City Hall, coated with Patina
• Watt's steam engine firebox due to superior heat dissipation. • Copper compounds in liquid form are used as a wood preservative, particularly in treating original portion of structures during restoration of damage due to dry rot. • Copper wires may be placed over non-conductive roofing materials to discourage the growth of moss. (Zinc may also be used for this purpose.) • Copper is used to prevent a building being directly struck by lightning. High above the roof, copper spikes (lightning rods) are connected to a very thick copper cable which leads to a large metal plate underneath the ground. The voltage is dispersed throughout the ground harmlessly, instead of destroying the main structure.[41]
Household products • Copper plumbing fittings and compression tubes. • Doorknobs and other fixtures in houses. • Roofing, guttering, and rainspouts on buildings. • In cookware, such as frying pans. • Some older flatware: (knives, forks, spoons) contains some copper if made from Electroplated Nickel silver (EPNS). • Sterling silver, if it is to be used in dinnerware, must contain a few percent copper. • Copper water heating cylinders • Copper Range Hoods • Copper Bath Tubs • Copper Counters
Old copper utensils in a Jerusalem restaurant
Copper • Copper Sinks • Copper slug tape
Coinage • As a component of coins, often as cupronickel alloy, or some form of brass or bronze. • Coins in the following countries all contain copper: European Union (Euro),[42] United States,[43] United Kingdom (sterling),[44] Australia[45] and New Zealand.[46] • U.S. Nickels are 75.0% copper by weight and only 25.0% nickel.[43]
Biomedical applications • As a biostatic surface in hospitals, and to line parts of ships to protect against barnacles and mussels, originally used pure, but superseded by Muntz metal. Bacteria will not grow on a copper surface because it is biostatic. Copper doorknobs are used by hospitals to reduce the transfer of disease, and Legionnaires' disease is suppressed by copper tubing in air-conditioning systems. • Copper(II) sulfate is used as a fungicide and as algae control in domestic lakes and ponds. It is used in gardening powders and sprays to kill mildew. • Copper-62-PTSM, a complex containing radioactive copper-62, is used as a positron emission tomography radiotracer for heart blood flow measurements. • Copper-64 can be used as a positron emission tomography radiotracer for medical imaging. When complexed with a chelate it can be used to treat cancer through radiation therapy.
Chemical applications • Compounds, such as Fehling's solution, have applications in chemistry. • As a component in ceramic glazes, and to color glass.
Other • Musical instruments, especially brass instruments and timpani. • Class D Fire Extinguisher, used in powder form to extinguish lithium fires by covering the burning metal and performing similar to a heat sink. • Textile fibers to create antimicrobial protective fabrics.[47] • Weaponry: • Small arms ammunition commonly uses copper as a jacketing material around the bullet core. • Copper is also commonly used as a case material, in the form of brass. • Copper is used as a liner in shaped-charge armor-piercing warheads. • Copper is frequently used in electroplating, usually as a base for other metals such as Nickel.
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Alloys Numerous copper alloys exist, many with important historical and contemporary uses. Speculum metal and bronze are alloys of copper and tin. Brass is an alloy of copper and zinc. Monel metal, also called cupronickel, is an alloy of copper and nickel. While the metal "bronze" usually refers to copper-tin alloys, it also is a generic term for any alloy of copper, such as aluminium bronze, silicon bronze, and manganese bronze. Copper is one of the most important constituents of carat silver and gold alloys and carat solders used in the jewelry industry, modifying the color, hardness and melting point of the resulting alloys.[48]
Compounds Common oxidation states of copper include the less stable copper(I) state, Cu+; and the more stable copper(II) state, Cu2+, which forms blue or blue-green salts and solutions. Under unusual conditions, a 3+ state and even an extremely rare 4+ state can be obtained. Using old nomenclature for the naming of salts, copper(I) is called cuprous, and copper(II) is cupric. In oxidation copper is mildly basic. Copper(II) carbonate is green from which arises the unique appearance of copper-clad roofs or domes on some buildings. Copper(II) sulfate forms a blue crystalline pentahydrate which is perhaps the most familiar copper compound in the laboratory. It is used as a fungicide, known as Bordeaux mixture. Copper (I) Oxide powder
There are two stable copper oxides, copper(II) oxide (CuO) and copper(I) oxide (Cu2O). Copper oxides are used to make yttrium barium copper oxide (YBa2Cu3O7-δ) or YBCO which forms the basis of many unconventional superconductors. • Copper(I) compounds: copper(I) chloride, copper(I) bromide, copper(I) iodide, copper(I) oxide. • Copper(II) compounds: copper(II) acetate, copper(II) carbonate, copper(II) chloride, copper(II) hydroxide, copper(II) nitrate, copper(II) oxide, copper(II) sulfate, copper(II) sulfide, copper(II) tetrafluoroborate, copper(II) triflate. • Copper(III) compounds, rare: potassium hexafluorocuprate (K3CuF6) • Copper(IV) compounds, extremely rare: caesium hexafluorocuprate (Cs2CuF6)
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Tests for copper(II) ion Adding an aqueous solution of sodium hydroxide will form a blue precipitate of copper(II) hydroxide. The ionic equation is: Cu2+ (aq) + 2 OH− (aq) → Cu(OH)2 (s) The full equation shows that the reaction is due to hydroxide ions deprotonating the hexaaquacopper(II) complex: [Cu(H2O)6]2+ (aq) + 2 OH−(aq) → Cu(H2O)4(OH)2 (s) + 2 H2O (l) Adding ammonium hydroxide (aqueous ammonia) causes the same precipitate to form. Upon adding excess ammonia, the precipitate dissolves, forming a deep blue ammonia complex, tetraamminecopper(II): Cu(H2O)4(OH)2 (s) + 4 NH3 (aq) → [Cu(H2O)2(NH3)4]2+ (aq) + 2 H2O (l) + 2 OH− (aq) A more delicate test than ammonia is potassium ferrocyanide, which gives a brown precipitate with copper salts.
Biological role Copper is essential in all plants and animals. The human body normally contains copper at a level of about 1.4 to 2.1 mg for each kg of body weight.[49] Copper is distributed widely in the body and occurs in liver, muscle and bone. Copper is transported in the bloodstream on a plasma protein called ceruloplasmin. When copper is first absorbed in the gut it is transported to the liver bound to albumin. Copper metabolism and excretion is controlled delivery of copper to the liver by ceruloplasmin, where it is excreted in bile. Copper is found in a variety of enzymes, including
Rich sources of Copper include oysters, beef or lamb liver, Brazil nuts, blackstrap molasses, cocoa, and black pepper. Good sources include lobster, nuts and sunflower seeds, green olives, avocados and wheat bran.
the copper centers of cytochrome c oxidase and the enzyme superoxide dismutase (containing copper and zinc). In addition to its enzymatic roles, copper is used for biological electron transport. The blue copper proteins that participate in electron transport include azurin and plastocyanin. The name "blue copper" comes from their intense blue color arising from a ligand-to-metal charge transfer (LMCT) absorption band around 600 nm. Most molluscs and some arthropods such as the horseshoe crab use the copper-containing pigment hemocyanin rather than iron-containing hemoglobin for oxygen transport, so their blood is blue when oxygenated rather than red.[50] It is believed that zinc and copper compete for absorption in the digestive tract so that a diet that is excessive in one of these minerals may result in a deficiency in the other. The RDA for copper in normal healthy adults is 0.9 mg/day. On the other hand, professional research on the subject recommends 3.0 mg/day.[51] Because of its role in facilitating iron uptake, copper deficiency can often produce anemia-like symptoms. In humans, the symptoms of Wilson's disease are caused by an accumulation of copper in body tissues.
Copper
Chronic copper depletion leads to abnormalities in Reference ranges for blood tests, comparing blood content of copper metabolism of fats, high (shown in light blue in middle) with other constituents. triglycerides, non-alcoholic steatohepatitis (NASH), fatty liver disease and poor melanin and dopamine synthesis causing depression and sunburn. Food rich in copper should be eaten away from any milk or egg proteins as they block absorption.
Toxicity Toxicity can occur from eating acidic food that has been cooked with copper cookware. Cirrhosis of the liver in children (Indian Childhood Cirrhosis) has been linked to boiling milk in copper cookware. The Merck Manual states that recent studies suggest that a genetic defect is associated with this cirrhosis.[52] Since copper is actively excreted by the normal body, chronic copper toxicosis in humans without a genetic defect in copper handling has not been demonstrated.[49] However, large amounts (gram quantities) of copper salts taken in suicide attempts have produced acute copper toxicity in normal humans. Equivalent amounts of copper salts (30 mg/kg) are toxic in animals[53]
Miscellaneous hazards The metal, when powdered, is a fire hazard. At concentrations higher than 1 mg/L, copper can stain clothes and items washed in water.
Recycling Copper is 100% recyclable without any loss of quality whether in a raw state or contained in a manufactured product. Copper is the third most recycled metal after iron and aluminum. It is estimated that 80% of the copper ever mined is still in use today.[54] Common grades of copper for recycling are: • Bare bright - very clean and pure copper wire usually 12 AWG or larger that has insulation and any tarnish removed • #1 copper - pipe with a new appearance and free of any foreign material • #2 copper - pipe with corrosion or foreign material and small gauge wire with no insulation Insulated wire is also commonly recycled.
See also • Copper mining and extraction: • Native copper • Companies: Anaconda Copper, Antofagasta PLC, Codelco, also Category:Copper mining companies • Copper extraction techniques, also Smelter • Peak copper • Electroplating • Pipe corrosion and damage: Cold water pitting of copper tube, Erosion corrosion of copper water tubes
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• Copper theft :Metal theft, also Operation Tremor
Further reading • "Copper: Technology & Competitiveness (Summary) Chapter 6: Copper Production Technology [55]". Office of Technology Assessment. 2005. http:/ / www. wws. princeton. edu/ ota/ ns20/ alpha_f. html. • Current Medicinal Chemistry, Volume 12, Number 10, May 2005, pp. 1161–1208(48) Metals, Toxicity and Oxidative Stress • William D. Callister (2003). Materials Science and Engineering: an Introduction, 6th Ed.. Table 6.1, p. 137: Wiley, New York. ISBN 0471736961. • Material: Copper (Cu), bulk [56], MEMS and Nanotechnology Clearinghouse. • Kim BE, Nevitt T, Thiele DJ (2008). "Mechanisms for copper acquisition, distribution and [57] [58] regulation ". Nat. Chem. Biol. 4: 176. doi:10.1038/nchembio.72 . PMID 18277979. http:/ / www. nature. com/ nchembio/ journal/ v4/ n3/ abs/ nchembio. 72. html. • Copper transport disorders [59]: an Instant insight from the Royal Society of Chemistry
External links • National Pollutant Inventory - Copper and compounds fact sheet
[60]
[61]
• Copper Resource Page. Includes 12 PDF files detailing the material properties of various kinds of copper, as well as various guides and tools for the copper industry. • The Copper Development Association [62] has an extensive site of properties and uses of copper; it also maintains a web site dedicated to brass, a copper alloy [63]. • The Third Millennium Online page on Copper [64] • The WebElements page on Copper [65] • Comprehensive Data on Copper [66]
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Article Sources and Contributors
Article Sources and Contributors Copper Source: http://en.wikipedia.org/w/index.php?oldid=305339601 Contributors: 007ketan, 07ireland07, 12dstring, 15236hb, 2D, 4timmy2turner0, 524, 786ace, A. B., A2Kafir, A314268, Abcdefghijklmnao, Abrech, Academic Challenger, Achim1999, Adam Johnston, Adam850, Adambro, Adamlala, Adashiel, Addshore, Adjusting, Ae-a, Agaudin,
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20
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21
Zinc
1
Zinc copper ← zinc → gallium ↑ Zn ↓ Cd
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
30Zn Periodic table
Appearance bluish pale gray
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
610
670
750
852
990
1179
Zinc
2
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1733.3 kJ·mol−1 3rd: 3833 kJ·mol−1Atomic radiusCovalent radiusVan der Waals radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusPoisson ratioMohs hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of zinc iso
N.A.
half-life
64
48.6%
64
65
syn
243.8 d
Zn Zn
27.9%
66
67
4.1%
67
68
18.8%
68
70
0.6%
70
72
syn
46.5 h
Zn Zn Zn Zn
DE (MeV)
DP
Zn is stable with 34 neutron
66
Zn
DM
ε
1.3519
65
γ
1.1155
-
β−
0.458
72
Cu
Zn is stable with 36 neutron Zn is stable with 37 neutron Zn is stable with 38 neutron Zn is stable with 40 neutron Ga
zinc, Zn, 30 transition metal12, 4, d65.38(4) g·mol−1 [Ar] 3d10 4s2 2, 8, 18, 2 (Image) solid 7.14 g·cm−3 6.57 g·cm−3 692.68 K,419.53 °C,787.15 °F 1180 K,907 °C,1665 °F 7.32 kJ·mol−1 123.6 kJ·mol−1 (25 °C) 25.470 J·mol−1·K−1+2, +1, 0 (amphoteric oxide) 1.65 (Pauling scale) 1st: 906.4 kJ·mol−1134 pm122±4 pm 139 pm hexagonal diamagnetic (20 °C) 59.0 nΩ·m (300 K) 116 W·m−1·K−1 (25 °C) 30.2 µm·m−1·K−1 (r.t.) (rolled) 3850 m·s−1 108 GPa 43 GPa 70 GPa 0.25 2.5 412 MPa 7440-66-6 Zinc (pronounced /ˈzɪŋk/, from German: Zink and also known as spelter) is a metallic chemical element with the symbol Zn and atomic number 30. It is a first-row transition metal in group 12 of the periodic table. Zinc is chemically similar to magnesium because its ion is of similar size and its only common oxidation state is +2. Zinc is the 24th most abundant element in the Earth's crust and has five stable isotopes. The most exploited zinc ore is sphalerite, or zinc sulfide; the largest exploitable deposits are found in Australia, Canada and the United States. Zinc production includes froth flotation of the ore, roasting and final extraction using electricity. Brass, which is an alloy of copper and zinc, has been used since at least the 10th century BC. Impure zinc metal was not produced in large scale until the 13th century in India, while the metal was unknown to Europe until the end of the 16th century. Alchemists burned zinc in air to form what they called "philosopher's wool" or "white snow." The element was probably named by the alchemist Paracelsus after the German word Zinke. German chemist Andreas Sigismund Marggraf is normally given credit for discovering pure metallic zinc in a 1746 experiment. Work by Luigi Galvani and Alessandro Volta uncovered the electrochemical properties of zinc by 1800. Corrosion-resistant zinc plating of steel is the major application for zinc. Other applications are in batteries and alloys, such as brass. A variety of zinc compounds are commonly used, such as zinc carbonate and zinc gluconate (as dietary supplements), zinc chloride (in deodorants), zinc pyrithione (anti-dandruff shampoos), zinc sulfide (in luminescent paints), and zinc methyl or
Zinc
3
zinc diethyl in the organic laboratory. Zinc is an essential mineral of "exceptional biologic and public health importance".[1] Zinc deficiency affects about 2 billion people in the developing world and is associated with many diseases.[2] In children it causes growth retardation, delayed sexual maturation, infection susceptibility, and diarrhea, contributing to the death of about 800,000 children worldwide per year.[1] Enzymes with a zinc atom in the reactive center are widespread in biochemistry, such as alcohol dehydrogenase in humans. Consumption of excess zinc can cause ataxia, lethargy and copper deficiency.
Characteristics Physical [3]
Zinc, also referred to in nonscientific contexts as spelter, is a bluish-white, lustrous, [4] diamagnetic metal, though most common commercial grades of the metal have a dull [5] finish. It is somewhat less dense than iron and has a hexagonal crystal structure.[6] The metal is hard and brittle at most temperatures but becomes malleable between 100 and 150 °C.[4] [5] Above 210 °C, the metal becomes brittle again and can be pulverized by beating.[7] Zinc is a fair conductor of electricity.[4] For a metal, zinc has relatively low melting (420 °C) and boiling points (900 °C).[8] Its melting point is the lowest of all the transition metals aside from mercury and cadmium.[8] Many alloys contain zinc, including brass, an alloy of zinc and copper. Other metals long known to form binary alloys with zinc are aluminium, antimony, bismuth, gold, iron, lead, mercury, silver, tin, magnesium, cobalt, nickel, tellurium and sodium.[9] While neither zinc nor zirconium are ferromagnetic, their alloy ZrZn2 exhibits ferromagnetism below 35 K.[4]
Occurrence Zinc makes up about 75 ppm (0.007%) of the Earth's crust, making it the 24th most abundant element there.[10] Soil contains 5–770 ppm of zinc with an average of 64 ppm.[10] Seawater has only 30 ppb zinc and the atmosphere contains 0.1–4 µg/m3.[10] The element is normally found in association with other base metals such as copper and lead in ores.[11] Zinc is a chalcophile ("sulfur loving"), meaning the element has a low affinity for oxygen and prefers to bond with sulfur in highly insoluble sulfides. Chalcophiles formed as the crust solidified under the reducing conditions of the early Earth's atmosphere.[12] Sphalerite, which is a form of zinc sulfide, is the most heavily mined zinc-containing ore because its concentrate contains 60–62% zinc.[11] Other minerals from which zinc is extracted include smithsonite (zinc carbonate), hemimorphite (zinc silicate), wurtzite (another zinc sulfide), and sometimes hydrozincite (basic zinc carbonate).[13] With the exception of wurtzite, all these other minerals were formed as a result of weathering processes on the primordial zinc sulfides.[12] Sphalerite (ZnS)
Zinc
4
World zinc resources total about 1.8 billion metric tons.[14] Nearly 200 million metric tons were economically viable in 2008; adding marginally economic and subeconomic reserves to that number, a total reserve base of 500 million metric tons has been identified.[14] Large deposits are in Australia, Canada and the United States.[12] At the current rate of consumption, these reserves are estimated to be depleted sometime between 2027 and 2055.[15] [16] About 346 teragrams (each teragram is equivalent to a megatonne) have been extracted throughout history to 2002, and one estimate found that about 109 teragrams of that remains in-use.[17]
Isotopes Five isotopes of zinc occur in nature. 64Zn is the most abundant isotope (48.63% natural abundance).[18] This isotope has such a long half-life, at ,[19] that its radioactivity can be ignored.[20] Similarly, 70Zn (0.6%), with a half life of is not usually considered to be radioactive. The other isotopes found in nature are 66Zn (28%), 67Zn (4%) and 68Zn (19%). Twenty-five radioisotopes have been characterized. 65Zn, which has a half-life of 243.66 days, is the most long-lived isotope, followed by 72Zn with a half-life of 46.5 hours.[18] Zinc has 10 nuclear isomers. 69mZn has the longest half-life, 13.76 h.[18] The superscript m indicates a metastable isotope. The nucleus of a metastable isotope is in an excited state and will return to the ground state by emitting a photon in the form of a gamma ray. 61Zn has three excitated states and 73Zn has two.[21] The isotopes 65Zn, 71Zn, 77 Zn and 78Zn each have only one excited state.[18] The most common decay mode of an isotope of zinc with a mass number lower than 64 is electron capture. The decay product resulting from electron capture is an isotope of copper.[18] n30Zn
+ e− →
n29Cu
The most common decay mode of an isotope of zinc with mass number higher than 64 is beta decay (β–), which produces an isotope of gallium.[18] n30Zn
→
n31Ga
+ e− + νe
Compounds and chemistry Reactivity Zinc has an electron configuration of [Ar]3d104s2 and is a member of the group 12 of the periodic table. It is a moderately reactive metal and strong reducing agent.[22] The surface of the pure metal tarnishes quickly, eventually forming a protective passivating layer of the basic zinc carbonate, Zn5(OH)6CO3, by reaction with atmospheric carbon dioxide.[23] This layer helps prevent further reaction with air and water. Zinc burns in air with a bright bluish-green flame, giving off fumes of zinc oxide.[24] Zinc reacts readily with acids, alkalis and other non-metals.[25] Extremely pure zinc reacts only slowly at room temperature with acids.[24] Strong acids, such as hydrochloric or sulphuric acid, can remove the passivating layer and subsequent reaction with water releases hydrogen gas.[24] The chemistry of zinc is dominated by the +2 oxidation state. When compounds in this oxidation state are formed the outer shell s electrons are lost, which yields a bare zinc ion with the electronic configuration [Ar]3d10.[26] This allows for the formation of four covalent
Zinc bonds by accepting four electron pairs and thus obeying the octet rule. The stereochemistry is therefore tetrahedral and the bonds may be described as being formed from sp3 hybrid orbitals on the zinc ion.[27] In aqueous solution an octahedral complex, [Zn(H2O)6]2+ is the predominant species.[28] The volatilization of zinc in combination with zinc chloride at temperatures above 285 °C indicates the formation of Zn2Cl2, a zinc compound with a +1 oxidation state.[24] No compounds of zinc in oxidation states other than +1 or +2 are known.[29] Calculations indicate that a zinc compound with the oxidation state of +4 is unlikely to exist.[30] Zinc chemistry is similar to the chemistry of the late first-row transition metals, nickel and copper though it has a filled d-shell, so its compounds are diamagnetic and mostly colorless.[31] The ionic radii of zinc and magnesium happen to be nearly identical. Because of this some of their salts have the same crystal structure[32] and in circumstances where ionic radius is a determining factor zinc and magnesium chemistries have much in common.[24] Otherwise there is little similarity. Zinc tends to form bonds with a greater degree of covalency and it forms much more stable complexes with N- and S- donors.[31] Complexes of zinc are mostly 4- or 6- coordinate although 5-coordinate complexes are known.[24]
Compounds Binary compounds of zinc are known for most of the metalloids and all the nonmetals except the noble gases. The oxide ZnO is a white powder that is nearly insoluble in neutral aqueous solutions, but is amphoteric, dissolving in both strong basic and acidic solutions.[24] The other chalcogenides (ZnS, ZnSe, and ZnTe) have varied applications in electronics and optics.[33] Pnictogenides (Zn3N2, Zn3P2, Zn3As2 and Zn3Sb2),[34] [35] the peroxide (ZnO2), the hydride (ZnH2), and the carbide (ZnC2) are also known.[36] Of Zinc chloride the four halides, ZnF2 has the most ionic character, whereas the others (ZnCl2, ZnBr2, and ZnI2) have relatively low melting points and are considered to have more covalent character.[37] In weak basic solutions containing Zn2+ ions, the hydroxide Zn(OH)2 forms as a white precipitate. In stronger alkaline solutions, this hydroxide is dissolved to form zincates ([Zn(OH)4]2-).[24] The nitrate Zn(NO3)2, chlorate Zn(ClO3)2, sulfate ZnSO4, phosphate Zn3(PO4)2, molybdate ZnMoO4, cyanide Zn(CN)2, arsenite Zn(AsO2)2, arsenate Zn(AsO4)2•8H2O and the chromate ZnCrO4 (one of the few colored zinc compounds) are a Basic zinc acetate few examples of other common inorganic compounds of zinc.[38] [39] One of the simplest examples of an organic compound of zinc is the acetate (Zn(O2CCH3)2). Organozinc compounds are those that contain zinc-carbon covalent bonds. Diethylzinc ((C2H5)2Zn) is a reagent in synthetic chemistry. It was first reported in 1848 from the reaction of zinc and ethyl iodide, and is the first compound known to contain a metal-carbon sigma bond.[40] Decamethyldizincocene contains a strong zinc-zinc bond at room
5
Zinc
6
temperature.[41]
History Ancient use Various isolated examples of the use of impure zinc in ancient times have been discovered. A possibly prehistoric statuette containing 87.5% zinc was found in a Dacian archaeological site in Transylvania (modern Romania).[42] Ornaments made of alloys that contain 80–90% zinc with lead, iron, antimony, and other metals making up the remainder, have been found that are 2500 years old.[11] The Berne zinc tablet is a votive plaque dating to Roman Gaul made of an alloy that is mostly zinc.[43] Zinc ores were used to make the zinc-copper alloy brass many centuries prior to the discovery of zinc as a separate element. Palestinian brass from the 14th to 10th centuries BC contains 23% zinc.[44] The Book of Genesis, written between the 10th and 5th centuries BC,[45] mentions Tubalcain as an "instructor in every artificer in brass and iron" (Genesis 4:22). Knowledge of how to produce brass spread to Ancient Greece by the 7th century BC but few varieties were made.[46]
Late Roman brass bucket the Hemmoorer Eimer from Warstade, Germany second to third century AD
The manufacture of brass was known to the Romans by about 30 BC.[47] They made brass by heating powdered calamine (zinc silicate or carbonate), charcoal and copper together in a crucible.[47] The resulting calamine brass was then either cast or hammered into shape and was used in weaponry.[48] Some coins struck by Romans in the Christian era are made of what is probably calamine brass.[49] In the West, impure zinc was known from antiquity to exist in the remnants in melting ovens, but it was usually discarded, as it was thought to be worthless.[50] References to medicinal uses of zinc are in the Charaka Samhita, which is believed to have been written as early as 300 BC in India.[51] The zinc mines and smelter of Zawar, near Udaipur in India, were active about 100 years before that[51] and produced an estimated million tonnes of metallic zinc and zinc oxide from the 12th to 16th centuries.[13] The Rasaratna Samuccaya, written in approximately the year 800, mentions two types of zinc-containing ores; one used for metal extraction and another used for medicinal purposes.[51]
Early studies and naming Zinc was distinctly recognized as a metal under the designation of Fasada in the medical Lexicon ascribed to the Hindu king Madanapala and written about the year 1374.[52] Smelting and extraction of impure zinc by reducing calamine with wool and other organic substances was accomplished in the 13th century in India.[4] [53] The Chinese did not learn of the technique until the 17th century.[53]
Zinc
7
Various alchemical symbols attributed to the element zinc
Alchemists burned zinc metal in air and collected the resulting zinc oxide on a condenser. Some alchemists called this zinc oxide lana philosophica, Latin for "philosopher's wool", because it collected in wooly tufts while others thought it looked like white snow and named it ninx album.[54]
The name of the metal was probably first used by Paracelsus, a Swiss-born German alchemist, who referred to the metal as "zincum" or "zinken" in his book Liber Mineralium II, in the 16th century.[53] [55] The word is probably derived from the German Zinke, and supposedly meant "tooth-like, pointed or jagged" (metallic zinc crystals have a needle-like appearance).[56] A second possibility is that the word is derived from the Persian word گنس seng meaning stone.[57] The metal was also called Indian tin, tutanego, calamine, and spinter.[11] German metallurgist Andreas Libavius received a quantity of what he called "calay" of Malabar from a cargo ship captured from the Portuguese in 1596.[58] Libavius described the properties of the sample, which may have been zinc. Zinc was regularly imported to Europe from the Orient in the 17th and early 18th centuries,[53] but was at times very expensive.[59]
Isolation of the pure element The isolation of metallic zinc in the West may have been achieved independently by several people. Postlewayt's Universal Dictionary, a contemporary source giving technological information in Europe, did not mention zinc before 1751 but the element was studied before then.[51] [60] Flemish metallurgist P.M. de Respour reported that he extracted metallic zinc from zinc oxide in 1668.[13] By the turn of the century, Étienne François Geoffroy described how zinc oxide condenses as yellow crystals on bars of iron placed above zinc ore being smelted.[13] In Britain, John Lane is said to have carried out experiments to smelt zinc, probably at Landore, prior to his bankruptcy in 1726.[61]
Credit for first isolating pure zinc is usually given to Andreas Sigismund Marggraf.
In 1738, William Champion patented in Great Britain a process to extract zinc from calamine in a vertical retort style smelter.[62] His technology was somewhat similar to that used at Zawar zinc mines in Rajasthan but there is no evidence that he visited the Orient.[63] Champion's process was used through 1851.[53] German chemist Andreas Marggraf normally gets credit for discovering pure metallic zinc even though Swedish chemist Anton von Swab distilled zinc from calamine four years before.[53] In his 1746 experiment, Marggraf heated a mixture of calamine and charcoal in a closed vessel without copper to obtain a metal.[50] This procedure became commercially practical by 1752.[64]
Zinc
8
Later work William Champion's brother, John, patented a process in 1758 for calcining zinc sulfide into an oxide usable in the retort process.[11] Prior to this only calamine could be used to produce zinc. In 1798, Johann Christian Ruberg improved on the smelting process by building the first horizontal retort smelter.[65] Jean-Jacques Daniel Dony built a different kind of horizontal zinc smelter in Belgium, which processed even more zinc.[53]
Galvanization was named for Luigi Galvani.
Italian doctor Luigi Galvani discovered in 1780 that connecting the spinal cord of a freshly dissected frog to an iron rail attached by a brass hook caused the frog's leg to twitch.[66] He incorrectly thought he had discovered an ability of nerves and muscles to create electricity and called the effect "animal electricity."[67] The galvanic cell and the process of galvanization were both named for Luigi Galvani and these discoveries paved the way for electrical batteries, galvanization and cathodic protection.[67]
Galvani's friend, Alessandro Volta, continued researching this effect and invented the Voltaic pile in 1800.[66] The basic unit of Volta's pile was a simplified galvanic cell, which is made of a plate of copper and a plate of zinc connected to each other externally and separated by an electrolyte. These were stacked in series to make the Voltaic cell, which in turn produced electricity by directing electrons from the zinc to the copper and allowing the zinc to corrode.[66] The non-magnetic character of zinc and its lack of color in solution delayed discovery of its importance to biochemistry and nutrition.[68] This changed in 1940 when carbonic anhydrase, an enzyme that scrubs carbon dioxide from blood, was shown to have zinc in its active site.[68] The digestive enzyme carboxypeptidase became the second known zinc-containing enzyme in 1955.[68]
Production Mining and processing Zinc is the fourth most common metal in use, trailing only iron, aluminium, and copper with an annual production of about 10 million tonnes.[] The world's largest zinc producer is Nyrstar, a merger of the Australian OZ Minerals and the Belgian Umicore.[69] About 70% of the world's zinc originates from mining, while the remaining 30% comes from recycling secondary zinc.[70] Commercially pure zinc is known as Special High Grade, often abbreviated SHG, and is 99.995% pure.[71]
Zinc
9
Worldwide, 95% of the zinc is mined from sulfidic ore deposits, in which sphalerite ZnS is nearly always mixed with the sulfides of copper, lead and iron.[73] There are zinc mines throughout the world, with the main mining areas being China, Australia and Peru.[] China produced over one-fourth of the global zinc output in 2006.[]
[72] Percentage of zinc output in 2006 by countries
Zinc metal is produced using extractive metallurgy.[74] After grinding the ore, froth flotation, which selectively separates minerals from gangue by taking advantage of differences in their hydrophobicity, is used to get an ore concentrate.[74] A final concentration of zinc of about 50% is reached by this process with the remainder of the concentrate being sulphur (32%), iron (13%), and SiO2 (5%).[74] Roasting converts the zinc sulfide concentrate produced during processing to zinc oxide:[73] 2 ZnS + 3 O2 → 2 ZnO + 2 SO2
Top 10 zinc producing countries in 2006 (full list) Rank
Country
tonnes
1
China (PRC)
2,600,000
2
Australia
1,338,000
3
Peru
1,201,794
4
United States
727,000
5
Canada
710,000
6
Mexico
480,000
7
Ireland
425,700
8
India
420,800
9
Kazakhstan
400,000
10
Sweden
192,400
The sulfur dioxide is used for the production of sulfuric acid, which is necessary for the leaching process. If deposits of zinc carbonate, zinc silicate or zinc spinel, like the Skorpion Deposit in Namibia are used for zinc production the roasting can be omitted.[75] For further processing two basic methods are used: pyrometallurgy or electrowinning. Pyrometallurgy processing reduces zinc oxide with carbon or carbon monoxide at 950 °C (1740 °F) into the metal, which is distilled as zinc vapor.[76] The zinc vapor is collected in a condenser.[73] The below set of equations demonstrate this process:[73] 2 ZnO + C → 2 Zn + CO2 2 ZnO + 2 CO → 2 Zn + 2 CO2 Electrowinning processing leaches zinc from the ore concentrate by sulfuric acid:[77]
Zinc
10 ZnO + H2SO4 → ZnSO4 + H2O
After this step electrolysis is used to produce zinc metal.[73]
Environmental impact The production for sulfidic zinc ores produces large amounts of sulfur dioxide and cadmium vapor. Smelter slag and other residues of process also contain significant amounts of heavy metals. About 1,100,000 tonnes of metallic zinc and 130,000 tonnes of lead were mined and smelted in the Belgian towns of La Calamine and Plombières between 1806 and 1882.[78] The dumps of the past mining operations leach significant amounts of zinc and cadmium, and, as a result, the sediments of the Geul River contain significant amounts of heavy metals.[78] About two thousand years ago emissions of zinc from mining and smelting totaled 10,000 tonnes a year. After increasing 10-fold from 1850, zinc emissions peaked at 3.4 megatonnes per year in the 1980s and declined to 2.7 megatonnes in the 1990s, although a 2005 study of the Arctic troposphere found that the concentrations there did not reflect the decline. Anthropogenic and natural emissions occur at a ratio of 20 to 1.[79] Levels of zinc in rivers flowing through industrial or mining areas can be as high as 20 ppm.[80] Effective sewage treatment greatly reduces this; treatment along the Rhine, for example, has decreased zinc levels to 50 ppb.[80] Concentrations of zinc as low as 2 ppm adversely affects the amount of oxygen that fish can carry in their blood.[81]
The zinc works at Lutana, is the largest exporter in Tasmania, generating 2.5% of the state's GDP. It produces over [82] 250000 tons of zinc per year. The Zinc works were historically responsible for high heavy metal levels in the [83] Derwent River
Soils contaminated with zinc through the mining of zinc-containing ores, refining, or where zinc-containing sludge is used as fertilizer, can contain several grams of zinc per kilogram (several ‰) of dry soil.[80] Levels of zinc in excess of 500 ppm in soil interfere with the ability of plants to absorb other essential metals, such as iron and manganese.[80] Zinc levels of 2000 ppm to 180,000 ppm (18%) have been recorded in some soil samples.[80]
Zinc
Applications Anti-corrosion and batteries The metal is most commonly used as an anti-corrosion agent.[84] Galvanization, which is the coating of iron or steel to protect the metals against corrosion, is the most familiar form of using zinc in this way. In 2006 in the United States, 56% or 773,000 tonnes of the zinc metal was used for galvanization,[85] while worldwide 47% was used for this purpose.[86] Zinc is more reactive than iron or steel and thus will attract almost all local oxidation until it completely Crystalline surface of a hot-dip corrodes away.[87] A protective surface layer of oxide galvanized handrail and carbonate (Zn5(OH)6(CO3)2) forms as the zinc corrodes.[88] This protection lasts even after the zinc layer is scratched but degrades through time as the zinc corrodes away.[88] The zinc is applied electrochemically or as molten zinc by hot-dip galvanizing or spraying.[10] Galvanization is used on chain-link fencing, guard rails, suspension bridges, lightposts, metal roofs, heat exchangers, and car bodies.[10] The relative reactivity of zinc and its ability to attract oxidation to itself also makes it a good sacrificial anode in cathodic protection. Cathodically protecting (CP) buried pipelines requires a solid piece of zinc to be connected by a conductor to a steel pipe.[88] Zinc acts as the anode (negative terminus) by slowly corroding away as it passes electric current to the steel pipeline.[88] [89] Zinc is also used to cathodically protect metals that are exposed to sea water from corrosion.[90] A zinc disc attached to a ship's iron rudder will slowly corrode while the rudder stays unattacked.[87] Other similar uses include a plug of zinc attached to a propeller or the metal protective guard for the keel of the ship. With an electrochemical potential of −0.7618 volts, zinc makes a good material for the negative terminus or anode in batteries.[91] Powdered zinc is used in this way in alkaline batteries and sheets of zinc metal form the cases for and act as anodes in zinc-carbon batteries.[92] [93]
11
Zinc
12
Alloys A widely used alloy of zinc is brass, in which copper is alloyed with anywhere from 3% to 45% zinc, depending upon the type of brass.[88] Brass is generally more ductile and stronger than copper and has superior corrosion resistance.[88] These properties make it useful in communication equipment, hardware, musical instruments, and water valves.[88]
Microstructure of cast brass at magnification 400X
Other widely used alloys that contain zinc include nickel silver, typewriter metal, soft and aluminum solder, and commercial bronze.[4] Zinc is also used in contemporary pipe organs as a substitute for the traditional lead/tin alloy in pipes.[94] Alloys of 85–88% zinc, 4–10% copper, and 2–8% aluminium find limited use in certain types of machine bearings. Zinc is the primary metal used in making American one cent coins since 1982.[95] The zinc core is coated with a thin layer of copper to give the impression of a copper coin. In 1994, 33200 tonnes (36600 short tons) of zinc were used to produce 13.6 billion pennies in the United
States.[96] Alloys of primarily zinc with small amounts of copper, aluminium, and magnesium are useful in die casting as well as spin casting, especially in the automotive, electrical, and hardware industries.[4] These alloys are marketed under the name Zamak.[97] An example of this is zinc aluminium. The low melting point together with the low viscosity of the alloy makes the production of small and intricate shapes possible. The low working temperature leads to rapid cooling of the cast products and therefore fast assembly is possible.[4] [86] [98] Another alloy, marketed under the name Prestal, contains 78% zinc and 22% aluminium and is reported to be nearly as strong as steel but as malleable as plastic.[4] This superplasticity of the alloy allows it to be molded using die casts made of ceramics and cement.[4] Similar alloys with the addition of a small amount of lead can be cold-rolled into sheets. An alloy of 96% zinc and 4% aluminium is used to make stamping dies for low production run applications for which ferrous metal dies would be too expensive.[99] In building facades, roofs or other applications in which zinc is used as sheet metal and for methods such as deep drawing, roll forming or bending, zinc alloys with titanium and copper are used.[100] Unalloyed zinc is too brittle for these kinds of manufacturing processes.[100] Cadmium zinc telluride (CZT) is a semiconductive alloy that can be divided into an array of small sensing devices.[101] These devices are similar to an integrated circuit and can detect the energy of incoming gamma ray photons.[101] When placed behind an absorbing mask, the CZT sensor array can also be used to determine the direction of the rays.[101] Zinc is used as the anode or fuel of the zinc-air battery/fuel cell providing the basis of the theorized zinc economy.[102] [103] [104]
Zinc
13
Other industrial uses Roughly one quarter of all zinc output, in the United States (2006), is consumed in the form of zinc compounds;[85] a variety of which are used industrially. Zinc oxide is widely used as a white pigment in paints, and as a catalyst in the manufacture of rubber.[10] It is also used as a heat disperser for the rubber and acts to protect its polymers from ultraviolet radiation (the same UV protection is conferred to plastics containing zinc oxide).[10] The semiconductor properties of zinc oxide make it useful in varistors and photocopying products.[105] The zinc zinc-oxide cycle is a two step thermochemical process based on zinc and zinc oxide for hydrogen production.[106]
Zinc oxide is used as a white pigment in paints.
Zinc chloride is often added to lumber as a fire retardant[107] and can be used as a wood preservative.[108] It is also used to make other chemicals.[107] Zinc methyl (Zn(CH3)2) is used in a number of organic syntheses.[109] Zinc sulfide (ZnS) is used in luminescent pigments such as on the hands of clocks, X-ray and television screens, and luminous paints.[110] Crystals of ZnS are used in lasers that operate in the mid-infrared part of the spectrum.[111] Zinc sulphate is a chemical in dyes and pigments.[107] Zinc pyrithione is used in antifouling paints.[112] Zinc powder is sometimes used as a propellant in model rockets.[113] When a compressed mixture of 70% zinc and 30% sulfur powder is ignited there is a violent chemical reaction.[113] This produces zinc sulfide, together with large amounts of hot gas, heat, and light.[113] Zinc sheet metal is used to make zinc bars.[114] Zinc has been proposed as a salting material for nuclear weapons (cobalt is another, better-known salting material).[115] A jacket of isotopically enriched Zn-64, irradiated by the intense high-energy neutron flux from an exploding thermonuclear weapon, would transmute into the radioactive isotope Zn-65 with a half-life of 244 days and produce massive gamma radiation, significantly increasing the radioactivity of the weapon's fallout for several days.[115] Such a weapon is not known to have ever been built, tested, or used.[115] Zn-65 is also used as a tracer to study how alloys that contain zinc wear out, or the path and the role of zinc in organisms.[116] Zinc dithiocarbamate complexes are used as agricultural fungicides, these include Zineb, Metiram, Propineb and Ziram.[117] Zinc naphthenate is used as wood preservative.[118]
Medicinal Zinc is included in most single tablet over-the-counter daily vitamin and mineral supplements.[119] It is believed to possess antioxidant properties, which protect against premature aging of the skin and muscles of the body, although studies differ as to its effectiveness.[120] Zinc also helps speed up the healing process after an injury.[120] Zinc gluconate glycine and zinc acetate are used in throat lozenges or tablets to reduce the duration and the severity of cold symptoms.[121] Preparations include zinc oxide, zinc acetate and zinc gluconate.[119]
Zinc
14 Zinc preparations can protect against sunburn in the summer and windburn in the winter.[47] Applied thinly to a baby's diaper area (perineum) with each diaper change, it can protect against diaper rash.[47] Zinc gluconate is one compound used for the delivery of zinc as a dietary supplement
The Age-Related Eye Disease Study determined that zinc can be part of an effective treatment for age-related macular degeneration.[122] Zinc supplementation is an effective treatment for acrodermatitis enteropathica, a genetic disorder affecting zinc absorption that was previously fatal to babies born with it.[47] Zinc lactate is used in toothpaste to prevent halitosis.[123] Zinc pyrithione is widely applied in shampoos because of its anti-dandruff function.[124] Zinc ions are effective antimicrobial agents even at low concentrations.[125] Gastroenteritis is strongly attenuated by ingestion of zinc, and this effect could be due to direct antimicrobial action of the zinc ions in the gastrointestinal tract, or to the absorption of the zinc and re-release from immune cells (all granulocytes secrete zinc), or both.[126] [127] [128] [129] [130] [131] [132]
Biological role Zinc is an essential trace element, necessary for plants,[79] , animals,[133] and microorganisms.[134] Zinc is found in nearly 100 specific enzymes[135] (other sources say 300), serves as structural ions in transcription factors and is stored and transferred in metallothioneins.[136] It is "typically the second most abundant transition metal in organisms" after iron and it is the only metal which appears in all enzyme classes.[79] The human body has 2-4 grams of zinc[137] distributed throughout the body. Most zinc is in the brain, muscle, bones, kidney, and liver, with the highest concentrations in the prostate and parts of the eye.[138] Semen is particularly rich in zinc, which is a key factor in prostate gland function and reproductive organ growth.[139] In humans zinc plays "ubiquitous biological roles".[1] It interacts with "a wide range of organic ligands"[1] and has a role in the metabolism of RNA and DNA, signal transduction, and gene expression. It also regulates apoptosis. A 2006 study estimated that about 10% of human proteins (2800) potentially bind zinc in addition to hundreds which transport and traffic zinc; a similar in silico study in the plant Arabidopsis thaliana found 2367 zinc-related proteins.[79] In the brain, zinc is stored in specific synaptic vesicles by glutamatergic neurons[140] and can "modulate brain excitability".[1] It plays a key role in synaptic plasticity and so in learning.[141] However it has been called "the brain's dark horse"[140] since it also can be a neurotoxin suggesting zinc homeostasis plays a critical role in normal functioning of the brain and central nervous system.[140]
Zinc
15
Enzymes Zinc is a good Lewis acid, making it a useful catalytic agent in hydroxylation and other enzymatic reactions.[135] The metal also has a flexible coordination geometry, which allows proteins using it to rapidly shift conformations to perform biological reactions.[142] Two examples of zinc-containing enzymes are carbonic anhydrase and carboxypeptidase, which are vital to the processes of carbon dioxide (CO2) regulation and digestion of proteins, respectively.[143] In vertebrate blood, carbonic anhydrase converts CO2 Ribbon diagram of human carbonic into bicarbonate and the same enzyme transforms the anhydrase II, with zinc atom visible in bicarbonate back into CO2 for exhalation through the the center [144] lungs. Without this enzyme, this conversion would occur about one million times slower[145] at the normal blood pH of 7 or would require a pH of 10 or more.[146] The non-related β-carbonic anhydrase is required in plants for leaf formation, the synthesis of indole acetic acid (auxin) and anaerobic respiration (alcoholic fermentation).[147] Carboxypeptidase cleaves peptide linkages during digestion of proteins. A coordinate covalent bond is formed between the terminal peptide and a C=O group attached to zinc, which gives the carbon a positive charge. This helps to create a hydrophobic pocket on the enzyme near the zinc, which attracts the non-polar part of the protein being digested.[143]
Other proteins Zinc serves a purely structural role in zinc fingers, twists and clusters.[148] Zinc fingers form parts of some transcription factors, which are proteins that recognize DNA base sequences during the replication and transcription of DNA. Each of the nine or ten Zn2+ ions in a zinc finger helps maintain the finger's structure by coordinately binding to four amino acids in the transcription factor.[145] The transcription factor wraps around the DNA helix and uses its fingers to accurately bind to the DNA sequence.
Zinc fingers help read DNA sequences
In blood plasma, zinc is bound to and transported by albumin (60 %, low-affinity) and transferrin (10%).[137] Since transferrin also transports iron, excessive iron reduces zinc absorption, and vice-versa. A similar reaction occurs with copper.[149] The concentration of zinc in blood plasma stays relatively constant regardless of zinc intake.[150] Cells in the salivary gland, prostate, immune system and intestine use zinc signaling as one way to communicate with other cells.[151]
Zinc may be held in metallothionein reserves within microorganisms or in the intestines or liver of animals.[152] Metallothionein in intestinal cells is capable of adjusting absorption of zinc by 15-40%.[153] However, inadequate or excessive zinc intake can be harmful; excess zinc particularly impairs copper absorption because metallothionein absorbs both metals.[154]
Zinc
16
Dietary intake In the U.S., the Recommended Dietary Allowance (RDA) is 8 mg/day for women and 11 mg/day for men.[155] Median intake in the U.S. around 2000 was 9 mg/day for women and 14 mg/day in men.[155] Red meats, especially beef, lamb and liver have some of the highest concentrations of zinc in food.[139] The concentration of zinc in plants varies based on levels of the element in soil. When there is adequate zinc in the soil, the food plants that contain the most zinc are wheat (germ and bran) and various seeds (sesame, poppy, alfalfa, celery, mustard).[156] Zinc is also found in beans, nuts, almonds, whole grains, pumpkin seeds, sunflower seeds and blackcurrant.[157] Soil conservation is needed to make sure that crop rotation will not deplete the zinc in soil.
Foods and spices that contain zinc
Other sources include fortified food and dietary supplements, which come in various forms. A 1998 review concluded that zinc oxide, one of the most common supplements in the United States, and zinc carbonate are nearly insoluble and poorly absorbed in the body.[158] This review cited studies which found low plasma zinc concentrations after zinc oxide and zinc carbonate were consumed compared with those seen after consumption of zinc acetate and sulfate salts.[158] However, harmful excessive supplementation is a problem among the relatively affluent, and should probably not exceed 20 mg/day in healthy people,[159] although the U.S. National Research Council set a Tolerable Upper Intake of 40 mg/day.[155] For fortification, however, a 2003 review recommended zinc oxide in cereals as cheap, stable, and as easily absorbed as more expensive forms.[160] A 2005 study found that various compounds of zinc, including oxide and sulfate, did not show statistically significant differences in absorption when added as fortificants to maize tortillas.[161]
Deficiency Zinc deficiency is usually due to insufficient dietary intake, but can be associated with malabsorption, acrodermatitis enteropathica, chronic liver disease, chronic renal disease, sickle cell disease, diabetes, malignancy, and other chronic illnesses.[2] Symptoms of mild zinc deficiency are diverse.[155] Clinical outcomes include depressed growth, diarrhea, impotence and delayed sexual maturation, alopecia, eye and skin lesions, impaired appetite, altered cognition, host defense properties, defects in carbohydrate utilization, and reproductive teratogenesis.[150] Mild zinc deficiency depresses immunity,[162] although so does excessive zinc.[137] Animals with a diet deficient in zinc require twice as much food to attain the same weight gain as animals given sufficient zinc.[110] Groups at risk for zinc deficiency include the elderly, vegetarians, and those with renal insufficiency. There is a paucity of adequate zinc biomarkers, and the most widely used indicator, plasma zinc, has poor sensitivity and specificity.[163] Diagnosing zinc deficiency is a persistent challenge.[1] Nearly 2 billion people in the developing world are deficient in zinc.[2] In children it causes an increase in infection and diarrhea, contributing to the death of about 800,000 children
Zinc worldwide per year.[1] The World Health Organization advocates zinc supplementation for severe malnutrition and diarrhea.[164] Zinc supplements help prevent disease and reduce mortality, especially among children with low birth weight or stunted growth.[164] However, zinc supplements should not be administered alone, since many in the developing world [165] have several deficiencies, and zinc interacts with other micronutrients. Zinc deficiency is plants' most common crop micronutrient deficiency; it is particularly common in high-pH soils. Zinc-deficient soil is cultivated in the cropland of about half of Turkey and India, a third of China, and most of Western Australia, and substantial responses to zinc fertilization have been reported in these areas.[79] Plants that grow in soils that are zinc-deficient are more susceptible to disease. Zinc is primarily added to the soil through the weathering of rocks, but humans have added zinc through fossil fuel combustion, mine waste, phosphate fertilizers, limestone, manure, sewage sludge, and particles from galvanized surfaces. Excess zinc is toxic to plants, although zinc toxicity is far less widespread.[79]
Precautions Toxicity Even though zinc is an essential requirement for good health, excess zinc can be harmful. Excessive absorption of zinc suppresses copper and iron absorption.[154] The free zinc ion in solution is highly toxic to plants, invertebrates, and even vertebrate fish.[166] The Free Ion Activity Model is well-established in the literature, and shows that just micromolar amounts of the free ion kills some organisms. A recent example showed 6 micromolar killing 93% of all Daphnia in water.[167] The free zinc ion is a powerful Lewis acid up to the point of being corrosive. Stomach acid contains hydrochloric acid, in which metallic zinc dissolves readily to give corrosive zinc chloride. Swallowing a post-1982 American one cent piece (97.5% zinc) can cause damage to the stomach lining due to the high solubility of the zinc ion in the acidic stomach.[168] There is evidence of induced copper deficiency at low intakes of 100–300 mg Zn/d; a recent trial had higher hospitalizations among elderly men taking 80 mg/day.[169] The USDA RDA is 15 mg Zn/d. Even lower levels, closer to the RDA, may interfere with the utilization of copper and iron or to adversely affect cholesterol.[154] Levels of zinc in excess of 500 ppm in soil interferes with the ability of plants to absorb other essential metals, such as iron and manganese.[80] There is also a condition called the zinc shakes or "zinc chills" that can be induced by the inhalation of freshly formed zinc oxide formed during the welding of galvanized materials.[110] The U.S. Food and Drug Administration (FDA) has stated that zinc damages nerve receptors in the nose, which can cause anosmia. Reports of anosmia were also observed in the 1930s when zinc preparations were used in a failed attempt to prevent polio infections. On June 16, 2009 the FDA said that consumers should stop using zinc-based intranasal cold products and ordered their removal from store shelves. The FDA said the loss of smell can be life-threatening because people with impaired smell cannot detect leaking gas or smoke and cannot tell if food has spoiled before they eat it.[170]
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Poisoning In 1982, the United States Mint began minting pennies coated in copper but made primarily of zinc. With the new zinc pennies, there is the potential for zinc toxicosis, which can be fatal. One reported case of chronic ingestion of 425 pennies (over 1 kg of zinc) resulted in death due to gastrointestinal bacterial and fungal sepsis, while another patient, who ingested 12 grams of zinc, only showed lethargy and ataxia (gross lack of coordination of muscle movements).[171] Several other cases have been reported of humans suffering zinc intoxication by the ingestion of zinc coins.[172] [173] Pennies and other small coins are sometimes ingested by dogs, resulting in the need for medical treatment to remove the foreign body. The zinc content of some coins can cause zinc toxicity, which is commonly fatal in dogs, where it causes a severe hemolytic anemia, and also liver or kidney damage; vomiting and diarrhea are possible symptoms.[174] Zinc is highly toxic in parrots and poisoning can often be fatal.[175] The consumption of fruit juices stored in galvanized cans has resulted in mass parrot poisonings with zinc.[47]
References [1] Hambidge, K. M. and Krebs, N. F. (2007). "Zinc deficiency: a special challenge". J. Nutr. 137: 1101. PMID 17374687. [2] Prasad, A. S. (2003). " Zinc deficiency (http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?tool=pmcentrez& artid=1125304)". British Medical Journal 326: 409. doi: 10.1136/bmj.326.7386.409 (http:/ / dx. doi. org/ 10. 1136/ bmj. 326. 7386. 409). PMID 12595353. [3] [4] [5] [6] [7]
" Spelter (http:/ / www. encyclo. co. uk/ define/ spelter)". Encyclo. . Retrieved 2009-08-01. CRC 2006, p. 4-41 Heiserman 1992, p. 123 Lehto 1968, p. 826 Scoffern, John (1861). The Useful Metals and Their Alloys (http:/ / books. google. com/ books?id=SSkKAAAAIAAJ). Houlston and Wright. pp. 591–603. . Retrieved 2009-04-06. [8] " Zinc Metal Properties (http:/ / www. galvanizeit. org/ aga/ designing-fabricating/ design-considerations/ zinc-metal-properties)". American Galvanizers Association. 2008. . Retrieved 2009-02-15. [9] Ingalls, Walter Renton (1902). Production and Properties of Zinc: A Treatise on the Occurrence and Distribution of Zinc Ore, the Commercial and Technical Conditions Affecting the Production of the Spelter, Its Chemical and Physical Properties and Uses in the Arts, Together with a Historical and Statistical Review of the Industry (http:/ / books. google. com/ books?id=RhNDAAAAIAAJ& pg=PA133& source=gbs_toc_r& cad=0_0#PPA127,M1). The Engineering and Mining Journal. pp. 142–146. . [10] Emsley 2001, p. 503 [11] Lehto 1968, p. 822 [12] Greenwood 1997, p. 1202 [13] Emsley 2001, p. 502 [14] Tolcin, A. C. (2009). " Mineral Commodity Summaries 2009: Zinc (http:/ / minerals. er. usgs. gov/ minerals/ pubs/ commodity/ zinc/ mcs-2009-zinc. pdf)" (PDF). United States Geological Survey. . Retrieved 2008-11-25. [15] Cohen, David (2007). "Earth audit". New Scientist 194: 8. doi: 10.1016/S0262-4079(07)61315-3 (http:/ / dx. doi. org/ 10. 1016/ S0262-4079(07)61315-3). [16] " Augsberg University Calculate When Our Materials Run Out (http:/ / www. idtechex. com/ products/ en/ articles/ 00000591. asp)". IDTechEx. 2007-06-04. . Retrieved 2008-12-09. [17] Gordon, R. B.; Bertram, M.; Graedel, T. E. (2006). " Metal stocks and sustainability (http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?tool=pmcentrez& artid=1360560)". Proceedings of the National Academy of Sciences 103: 1209. doi: 10.1073/pnas.0509498103 (http:/ / dx. doi. org/ 10. 1073/ pnas. 0509498103). PMID 16432205. [18] NNDC contributors (2008). " Chart of Nuclides (http:/ / www. nndc. bnl. gov/ chart/ )". Upton (NY): National Nuclear Data Center, Brookhaven National Laboratory. . Retrieved 2008-09-13. [19] CRC 2006, p. 11-70 [20] NASA contributors. " Five-Year Wilkinson Microwave Anisotropy Probe (WMAP) Observations: Data Processing, Sky Maps, and Basic Results (http:/ / lambda. gsfc. nasa. gov/ product/ map/ dr3/ pub_papers/ fiveyear/ basic_results/ wmap5basic. pdf)" (PDF). nasa.gov. . Retrieved 2008-03-06.
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Zinc [21] Audi, Georges (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3–128. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [22] CRC 2006, pp. 8-29 [23] Porter, Frank C. (1994). Corrosion Resistance of Zinc and Zinc Alloys. CRC Press. p. 121. ISBN 0-8247-9213-0. [24] Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils; (1985). "Zink" (in German). Lehrbuch der Anorganischen Chemie (91–100 ed.). Walter de Gruyter. pp. 1034–1041. ISBN 3110075113. [25] Hinds, John Iredelle Dillard (1908). Inorganic Chemistry: With the Elements of Physical and Theoretical Chemistry (http:/ / books. google. com/ books?id=xMUMAAAAYAAJ) (2nd ed.). New York: John Wiley & Sons. pp. 506–508. . [26] Ritchie, Rob (2004). Chemistry (http:/ / books. google. com/ books?id=idT9j6406gsC) (2nd ed.). Letts and Lonsdale. p. 71. ISBN 1-84315-438-2. . [27] Jaffe, Howard W. (1996). Crystal Chemistry and Refractivity (http:/ / books. google. com/ books?id=lbxUYhC5YEEC). Courier Dover Publications. p. 31. ISBN 0-486-69173-X. . [28] Burgess, John (1978). Metal ions in solution. New York: Ellis Horwood. p. 147. [29] Brady, James E.; Humiston, Gerard E.; Heikkinen, Henry (1983). General Chemistry: Principles and Structure (3rd ed.). John Wiley & Sons. p. 671. ISBN 0-471-86739-X. [30] Kaupp M.; Dolg M.; Stoll H.; Von Schnering H. G. (1994). " Oxidation state +IV in group 12 chemistry. Ab initio study of zinc(IV), cadmium(IV), and mercury(IV) fluorides (http:/ / pubs. acs. org/ doi/ pdf/ 10. 1021/ ic00088a012)". Inorganic chemistry 33: 2122. doi: 10.1021/ic00088a012 (http:/ / dx. doi. org/ 10. 1021/ ic00088a012). . [31] Greenwood 1997, p. 1206 [32] CRC 2006, pp. 12-11–12 [33] " Zinc Sulfide (http:/ / www. americanelements. com/ znsu. html)". American Elements. . Retrieved 2009-02-03. [34] Grolier contributors (1994). Academic American Encyclopedia (http:/ / books. google. com/ books?id=YgI4E7w5JI8C). Danbury, Connecticut: Grolier Inc.. p. 202. ISBN 0-7172-2053-2. . Retrieved 2007-11-01. [35] " Zinc Phosphide (http:/ / www. americanelements. com/ znp. html)". American Elements. . Retrieved 2009-02-03. [36] Shulzhenko, A. A.; Ignatyeva, I. Yu.; Osipov, A. S.; Smirnova, T. I. (2000). "Peculiarities of interaction in the Zn–C system under high pressures and temperatures". Diamond and Related Materials 9: 129. doi: 10.1016/S0925-9635(99)00231-9 (http:/ / dx. doi. org/ 10. 1016/ S0925-9635(99)00231-9). [37] Greenwood 1997, p. 1211 [38] Rasmussen, J. K.; Heilmann, S. M. (1990). " In situ Cyanosilylation of Carbonyl Compounds: O-Trimethylsilyl-4-Methoxymandelonitrile (http:/ / www. orgsyn. org/ orgsyn/ prep. asp?prep=cv7p0521)". Organic Syntheses, Collected Volume] 7: 521. . Retrieved 2009-04-06. [39] Perry, D. L. (1995). Handbook of Inorganic Compounds. CRC Press. pp. 448–458. ISBN 0849386713. [40] Frankland, E. (1850). "On the isolation of the organic radicals". Quarterly Journal of the Chemical Society 2: 263. doi: 10.1039/QJ8500200263 (http:/ / dx. doi. org/ 10. 1039/ QJ8500200263). [41] Resa, I.; Carmona, E.; Gutierrez-Puebla, E.; Monge, A. (2004). "Decamethyldizincocene, a Stable Compound of Zn(I) with a Zn-Zn Bond". Science 304: 1136. doi: 10.1126/science.1101356 (http:/ / dx. doi. org/ 10. 1126/ science. 1101356). PMID 15326350. [42] Weeks 1933, p. 20 [43] Rehren, Th. (1996). "A Roman zinc tablet from Bern, Switzerland: Reconstruction of the Manufacture". in S. Demirci et al.. The Proceedings of the 29th International Symposium on Archaeometry. 94. Archaeometry. pp. 35–45. [44] Greenwood 1997, p. 1201 [45] Lerner, Gerda (2003). " Religion and the Creation of Feminist Consciousness (http:/ / www. hds. harvard. edu/ news/ bulletin/ articles/ lerner. html)". Harvard Divinity Bulletin 23 (1). . Retrieved 2009-04-06. [46] Craddock, Paul T. (1978). "The composition of copper alloys used by the Greek, Etruscan and Roman civilizations. The origins and early use of brass". Journal of Archaeological Science 5: 1. doi: 10.1016/0305-4403(78)90015-8 (http:/ / dx. doi. org/ 10. 1016/ 0305-4403(78)90015-8). [47] Emsley 2001, p. 501 [48] " How is zinc made? (http:/ / www. answers. com/ zinc)". How Products are Made. The Gale Group. 2002. . Retrieved 2009-02-21. [49] Chambers 1901, p. 799 [50] Weeks 1933, p. 21
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Zinc [51] Craddock, P. T.; Gurjar L. K.; Hegde K. T. M. (1983). " Zinc production in medieval India (http:/ / www. jstor. org/ pss/ 124653)". World Archaeology 15 (2): 211. . [52] Ray, Prafulla Chandra (1903). A History of Hindu Chemistry from the Earliest Times to the Middle of the Sixteenth Century, A.D.: With Sanskrit Texts, Variants, Translation and Illustrations (http:/ / books. google. com/ books?id=DL1HAAAAIAAJ& printsec=titlepage& source=gbs_summary_r& cad=0). 1 (2nd ed.). The Bengal Chemical & Pharmaceutical Works, Ltd.. pp. 157–158. . (public domain text) [53] Habashi, Fathi. " Discovering the 8th Metal (http:/ / www. iza. com/ Documents/ Communications/ Publications/ History. pdf)" (PDF). International Zinc Association (IZA). . Retrieved 2008-12-13. [54] Arny, Henry Vinecome (1917). Principles of Pharmacy (http:/ / books. google. com/ books?id=gRNKAAAAMAAJ) (2nd ed.). W. B. Saunders company. p. 483. . [55] Hoover, Herbert Clark (2003). Georgius Agricola de Re Metallica. Kessinger Publishing. p. 409. ISBN 0-7661-3197-1. [56] Gerhartz, Wolfgang; et al. (1996). Ullmann's Encyclopedia of Industrial Chemistry (5th ed.). VHC. p. 509. ISBN 3527201009. [57] Fathi Habashi (1997). Handbook of Extractive Metallurgy. Wiley-VHC. p. 642. ISBN 3-527-28792-2. [58] Lach, Donald F. (1994). " Technology and the Natural Sciences (http:/ / books. google. com/ books?id=N0xD7BYXv_YC& pg=PA426)". Asia in the Making of Europe. University of Chicago Press. p. 426. ISBN 0226467341. . [59] An East India Company ship carrying a cargo of nearly pure zinc metal from the Orient sank off the coast Sweden in 1745.(Emsley 2001, p. 502) [60] Lynn Willies, P. T. Craddock, L. J. Gurjar and K. T. M. Hedge World Archaeology (1984). Ancient Lead and Zinc Mining in Rajasthan, India (http:/ / www. jstor. org/ stable/ 124574). 16. p. 222. . [61] Roberts, R. O. (1951). "Dr John Lane and the foundation of the non-ferrous metal industry in the Swansea valley". Gower (Gower Society) (4): 19. [62] Comyns, Alan E. (2007). Encyclopedic Dictionary of Named Processes in Chemical Technology (http:/ / books. google. com/ books?id=Jlq-ckWvQSQC) (3rd ed.). CRC Press. p. 71. ISBN 0-8493-9163-6. . [63] Jenkins, Rhys (1945–7). "The Zinc Industry in England: the early years up to 1850". Transactions of the Newcomen Society 25: 41–52. [64] Heiserman 1992, p. 122 [65] Gray (2005). Zinc. Marshall Cavendish. p. 8. ISBN 0-7614-1922-5. [66] Warren, Neville G. (2000). Excel Preliminary Physics (http:/ / books. google. com/ books?id=eL9Xn6nQ6XQC& printsec=frontcover& source=gbs_summary_r& cad=0). Pascal Press. p. 47. ISBN 1-74020-085-3. . [67] " Galvanic Cell (http:/ / books. google. com/ books?id=gV1MAAAAMAAJ& pg=PA80)". The New International Encyclopaedia. Dodd, Mead and Company. 1903. p. 80. . [68] Cotton 1999, p. 626 [69] Pearson, Madelene; Ann, Tan Hwee (December 12, 2006). " Zinifex and Umicore to create largest zinc producer (http:/ / www. iht. com/ articles/ 2006/ 12/ 12/ bloomberg/ sxzini. php)". Bloomberg News (International Herald Tribune). . Retrieved 2008-11-24. [70] " Zinc Recycling (http:/ / www. zincworld. org/ recycling. html)". International Zinc Association. . Retrieved 2008-11-28. [71] " Special High Grade Zinc (SHG) 99.995% (http:/ / nyrstar. com/ nyrstar/ en/ products/ zinccongalvanising/ techdownloads/ shg_budel. pdf)" (PDF). Nyrstar. 2008. . Retrieved 2008-12-01. [72] Jasinski, Stephen M.. " Mineral Commodity Summaries 2007: Zinc (http:/ / minerals. er. usgs. gov/ minerals/ pubs/ commodity/ zinc/ mcs-2008-zinc. pdf)" (PDF). United States Geological Survey. . Retrieved 2008-11-25. [73] Porter, Frank C. (1991). Zinc Handbook (http:/ / books. google. com/ books?& id=laACw9i0D_wC). CRC Press. ISBN 978-0-8247-8340-2. . [74] Rosenqvist, Terkel (1922). Principles of Extractive Metallurgy (2 ed.). Tapir Academic Press. pp. 7, 16, 186. ISBN 82-519-1922-3. [75] Borg, Gregor; Kärner, Katrin; Buxton, Mike; Armstrong, Richard; van der Merwe, Schalk W. (2003). "Geology of the Skorpion Supergene Zinc Deposit, Southern Namibia". Economic Geology 98: 749. doi: 10.2113/98.4.749 (http:/ / dx. doi. org/ 10. 2113/ 98. 4. 749). [76] Bodsworth, Colin (1994). The Extraction and Refining of Metals. CRC Press. p. 148. ISBN 0-8493-4433-6. [77] Gupta, C. K.; Mukherjee, T. K. (1990). Hydrometallurgy in Extraction Processes. CRC Press. p. 62. ISBN 0-8493-6804-9. [78] Kucha, H.; Martens, A.; Ottenburgs, R.; De Vos, W.; Viaene, W. (1996). "Primary minerals of Zn-Pb mining and metallurgical dumps and their environmental behavior at Plombières, Belgium". Environmental Geology 27: 1. doi: 10.1007/BF00770598 (http:/ / dx. doi. org/ 10. 1007/ BF00770598).
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Zinc [79] Broadley, M. R.; White, P. J.; Hammond, J. P.; Zelko I.; Lux A. (2007). "Zinc in plants". New Phytologist 173: 677. doi: 10.1111/j.1469-8137.2007.01996.x (http:/ / dx. doi. org/ 10. 1111/ j. 1469-8137. 2007. 01996. x). PMID 17286818. [80] Emsley 2001, p. 504 [81] Heath, Alan G. (1995). Water pollution and fish physiology (http:/ / books. google. com/ books?id=5NPVTuBtGF4C). Boca Raton, Florida: CRC Press. p. 57. ISBN 0-87371-632-9. . [82] " "The Zinc Works" (http:/ / www. tchange. com. au/ resources/ zinifex_smelter. html)". TChange. . Retrieved 2009-07-11. [83] " Derwent Estuary - Water Quality Improvement Plan for Heavy Metals (http:/ / www. derwentestuary. org. au/ file. php?id=193)". Derwent Estuary Program. June 2007. . Retrieved 2009-07-11. [84] Greenwood 1997, p. 1203 [85] Tolcin, Amy C.. " Mineral Yearbook 2006: Zinc (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ zinc/ zinc_mcs06. pdf)" (PDF). United States Geological Survey. . Retrieved 2009-04-06. [86] Panagapko, Doug (2006). " Zinc (http:/ / info. wlu. ca/ ~wwwgeog/ special/ vgt/ English/ can_mod2/ unit7. htm)". Natural Resources Canada. . Retrieved 2008-12-12. [87] Stwertka 1998, p. 99 [88] Lehto 1968, p. 829 [89] Electric current will naturally flow between zinc and steel but larger pipeline systems require a rectifier that adds an induced DC electric current to the CP system. [90] Bounoughaz, M.; Salhi, E.; Benzine, K.; Ghali E.; Dalard F. (2003). "A comparative study of the electrochemical behaviour of Algerian zinc and a zinc from a commercial sacrificial anode". Journal of Materials Science 38: 1139. doi: 10.1023/A:1022824813564 (http:/ / dx. doi. org/ 10. 1023/ A:1022824813564). [91] CRC 2006, p. 8–25 [92] Besenhard, Jürgen O. (1999) (PDF). Handbook of Battery Materials (http:/ / www. ulb. tu-darmstadt. de/ tocs/ 60178752. pdf). Wiley-VCH. ISBN 3527294694. . Retrieved 2008-10-08. [93] Wiaux, J. -P.; Waefler, J. -P. (1995). "Recycling zinc batteries: an economical challenge in consumer waste management". Journal of Power Sources 57: 61. doi: 10.1016/0378-7753(95)02242-2 (http:/ / dx. doi. org/ 10. 1016/ 0378-7753(95)02242-2). [94] Bush, Douglas Earl; Kassel, Richard (2006). The Organ: An Encyclopedia (http:/ / books. google. com/ books?id=cgDJaeFFUPoC). Routledge. p. 679. ISBN 978-0-415-94174-7. . [95] " Coin Specifications (http:/ / www. usmint. gov/ about_the_mint/ ?action=coin_specifications)". United States Mint. . Retrieved 2008-10-08. [96] Jasinski, Stephen M.. " Mineral Yearbook 1994: Zinc (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ zinc/ 720494. pdf)" (PDF). United States Geological Survey. . Retrieved 2008-11-13. [97] Eastern Alloys contributors. " Diecasting Alloys (http:/ / www. eazall. com/ diecastalloys. aspx)". Maybrook, NY: Eastern Alloys. . Retrieved 2009-01-19. [98] Apelian, D.; Paliwal, M.; Herrschaft, D. C. (1981). "Casting with Zinc Alloys". Journal of Metals 33: 12–19. [99] Samans, Carl Hubert (1949). Engineering Metals and Their Alloys. Macmillan Co.. [100] Porter, Frank (1994). " Wrought Zinc (http:/ / books. google. com/ books?id=C-pAiedmqp8C)". Corrosion Resistance of Zinc and Zinc Alloys. CRC Press. pp. 6–7. ISBN 9780824792138. . [101] Katz, Johnathan I. (2002). The Biggest Bangs. Oxford University Press. p. 18. ISBN 0-19-514570-4. [102] Culter, T. (1996). "A design guide for rechargeable zinc-air battery technology". Southcon/96. Conference Record: 616. doi: 10.1109/SOUTHC.1996.535134 (http:/ / dx. doi. org/ 10. 1109/ SOUTHC. 1996. 535134). [103] Whartman, Jonathan. " Zinc Air Battery-Battery Hybrid for Powering Electric Scooters and Electric Buses (http:/ / www. electric-fuel. com/ evtech/ papers/ paper11-1-98. pdf)" (PDF). The 15th International Electric Vehicle Symposium. . Retrieved 2008-10-08. [104] Cooper, J. F; Fleming, D.; Hargrove, D.; Koopman, R.; Peterman, K.. " A refuelable zinc/air battery for fleet electric vehicle propulsion (http:/ / www. osti. gov/ energycitations/ product. biblio. jsp?osti_id=82465)". Society of Automotive Engineers future transportation technology conference and exposition. . Retrieved 2008-10-08. [105] Zhang, Xiaoge Gregory (1996). Corrosion and Electrochemistry of Zinc (http:/ / books. google. com/ books?id=Qmf4VsriAtMC). Springer. p. 93. ISBN 0-306-45334-7. . [106] Weimer, Al (2006-05-17). " Development of Solar-powered Thermochemical Production of Hydrogen from Water (http:/ / www. hydrogen. energy. gov/ pdfs/ review06/ pd_10_weimer. pdf)" (PDF). U.S. Department of Energy. . Retrieved 2009-01-10. [107] Heiserman 1992, p. 124 [108] Blew, Joseph Oscar (1953). Wood preservatives (http:/ / hdl. handle. net/ 1957/ 816). Department of Agriculture, Forest Service, Forest Products Laboratory. . [109] Frankland, Edward (1849). "Notiz über eine neue Reihe organischer Körper, welche Metalle, Phosphor u. s. w. enthalten" (in German). Liebig's Annalen der Chemie und Pharmacie 71: 213. doi: 10.1002/jlac.18490710206
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(http:/ / dx. doi. org/ 10. 1002/ jlac. 18490710206). [110] CRC 2006, p. 4-42 [111] Paschotta, Rüdiger (2008). Encyclopedia of Laser Physics and Technology (http:/ / books. google. com/ books?id=BN026ye2fJAC). Wiley-VCH. p. 798. ISBN 3-527-40828-2. . [112] Konstantinou, I. K.; Albanis, T. A. (2004). "Worldwide occurrence and effects of antifouling paint booster biocides in the aquatic environment: a review". Environment International 30: 235. doi: 10.1016/S0160-4120(03)00176-4 (http:/ / dx. doi. org/ 10. 1016/ S0160-4120(03)00176-4). [113] Boudreaux, Kevin A.. " Zinc + Sulfur (http:/ / www. angelo. edu/ faculty/ kboudrea/ demos/ zinc_sulfur/ zinc_sulfur. htm)". Angelo State University. . Retrieved 2008-10-08. [114] " Technical Information (http:/ / www. zinccounters. co. uk/ html/ tech/ tech. htm)". Zinc Counters. 2008. . Retrieved 2008-11-29. [115] Win, David Tin; Masum, Al (2003). " Weapons of Mass Destruction (http:/ / www. journal. au. edu/ au_techno/ 2003/ apr2003/ aujt6-4_article07. pdf)" (PDF). Assumption University Journal of Technology (Assumption University) 6 (4): 199. . Retrieved 2009-04-06. [116] David E. Newton (1999). Chemical Elements: From Carbon to Krypton (http:/ / www. encyclopedia. com/ doc/ 1G2-3427000114. html). U. X. L. /Gale. ISBN 0-7876-2846-8. . Retrieved 2009-04-06. [117] Ullmann's Agrochemicals (http:/ / books. google. com/ books?id=cItuoO9zSjkC& pg=PA591). Wiley-Vch (COR). 2007. pp. 591–592. ISBN 3527316043. . [118] Walker, J. C. F. (2006). Primary Wood Processing: Principles and Practice. Springer. p. 317. ISBN 1402043929. [119] DiSilvestro, Robert A. (2004). Handbook of Minerals as Nutritional Supplements. CRC Press. pp. 135, 155. [120] Milbury, Paul E.; Richer, Alice C. (2008). Understanding the Antioxidant Controversy: Scrutinizing the "fountain of Youth". Greenwood Publishing Group. p. 99. ISBN 0-275-99376-0. [121] Ananda S., Prasad; Fitzgerald, James T.; Bao, Bin; Beck, Frances W.J.; Chandrasekar, Pranatharthi H. (2000). " Duration of Symptoms and Plasma Cytokine Levels in Patients with the Common Cold Treated with Zinc Acetate: A Randomized, Double-Blind, Placebo-Controlled Trial (http:/ / www. annals. org/ cgi/ reprint/ 133/ 4/ 245. pdf)" (PDF). Annals of Internal Medicine 133 (4): 245. . [122] Age-Related Eye Disease Study Research Group (2001). " A Randomized, Placebo-Controlled, Clinical Trial of High-Dose Supplementation With Vitamins C and E, Beta Carotene, and Zinc for Age-Related Macular Degeneration and Vision Loss (http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?tool=pubmed& pubmedid=11594942)". Arch Ophthalmology 119 (10): 1417. PMID 11594942. . [123] Roldán, S.; Winkel, E. G.; Herrera, D.; Sanz, M.; Van Winkelhoff, A. J. (2003). "The effects of a new mouthrinse containing chlorhexidine, cetylpyridinium chloride and zinc lactate on the microflora of oral halitosis patients: a dual-centre, double-blind placebo-controlled study". Journal of Clinical Periodontology 30: 427. doi: 10.1034/j.1600-051X.2003.20004.x (http:/ / dx. doi. org/ 10. 1034/ j. 1600-051X. 2003. 20004. x). [124] Marks, R.; Pearse, A. D.; Walker, A. P. (1985). "The effects of a shampoo containing zinc pyrithione on the control of dandruff". British Journal of Dermatology 112: 415. doi: 10.1111/j.1365-2133.1985.tb02314.x (http:/ / dx. doi. org/ 10. 1111/ j. 1365-2133. 1985. tb02314. x). [125] McCarthy, T J; Zeelie, J J: Krause, D J (1992 Feb). "The antimicrobial action of zinc ion/antioxidant combinations.". Clinical Pharmacology & Therapeutics (American Society for Clinical Pharmacology and Therapeutics) 17 (1): 5. [126] Aydemir, T. B.; Blanchard, R. K.; Cousins, R. J. (2006). "Zinc Supplementation of Young Men Alters Metallothionein, Zinc Transporter, and Cytokine Gene Expression in Leucocyte Populations". PNAS 103: 1699. doi: 10.1073/pnas.0510407103 (http:/ / dx. doi. org/ 10. 1073/ pnas. 0510407103). [127] Valko, M.; Morris, H.; Cronin, M. T. D. (2005). "Metals, Toxicity and Oxidative stress". Current Medicinal Chemistry 12: 1161. doi: 10.2174/0929867053764635 (http:/ / dx. doi. org/ 10. 2174/ 0929867053764635). [128] In clinical trials, both zinc gluconate and zinc gluconate glycine (the formulation used in lozenges) have been shown to shorten the duration of symptoms of the common cold.
Godfrey, J. C.; Godfrey, N. J.; Novick, S. G. (1996). "Zinc for treating the common cold: Review of all clinical trials since 1984". Alternative Therapies in Health and Medicine. PMID 8942045. [130] The amount of glycine can vary from two to twenty moles per mole of zinc gluconate. One review of the research found that out of nine controlled experiments using zinc lozenges, the results were positive in four studies, and no better than placebo in five.
Hulisz, Darrell T.. " Zinc and the Common Cold: What Pharmacists Need to Know (http:/ / www. uspharmacist. com/ oldformat. asp?url=newlook/ files/ alte/ feat2. htm)". uspharmacist.com. . Retrieved 2008-11-28.
Zinc [132] This review also suggested that the research is characterized by methodological problems, including differences in the dosage amount used, and the use of self-report data. The evidence suggests that zinc supplements may be most effective if they are taken at the first sign of cold symptoms. [133] Prasad A. S. (2008). " Zinc in human health: effect of zinc on immune cells (http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?tool=pmcentrez& artid=2277319)". Mol. Med. 14: 353. doi: 10.2119/2008-00033.Prasad (http:/ / dx. doi. org/ 10. 2119/ 2008-00033. Prasad). PMID 18385818. [134] Zinc's role in microorganisms is particularly reviewed in: Sugarman B (1983). "Zinc and infection". Review of Infectious Diseases 5: 137. PMID 6338570. [135] NRC 2000, p. 443 [136] Cotton 1999, pp. 625–629 [137] Rink, L.; Gabriel P. (2000). "Zinc and the immune system". Proc Nutr Soc 59: 541. doi: 10.1017/S0029665100000781 (http:/ / dx. doi. org/ 10. 1017/ S0029665100000781). PMID 11115789. [138] Wapnir, Raul A. (1990). Protein Nutrition and Mineral Absorption (http:/ / books. google. com/ books?id=qfKdaCoZS18C). Boca Raton, Florida: CRC Press. ISBN 0849352274. . [139] Berdanier, Carolyn D.; Dwyer, Johanna T.; Feldman, Elaine B. (2007). Handbook of Nutrition and Food (http:/ / books. google. com/ books?id=PJpieIePsmUC). Boca Raton, Florida: CRC Press. ISBN 0849392187. . [140] Bitanihirwe BK, Cunningham MG (2009). "Zinc: The brain's dark horse". Synapse 63: 1029. PMID 19623531. [141] Nakashima AS, Dyck RH (2009). "Zinc and cortical plasticity". Brain Res Rev 59: 347. doi: 10.1016/j.brainresrev.2008.10.003 (http:/ / dx. doi. org/ 10. 1016/ j. brainresrev. 2008. 10. 003). PMID 19026685. [142] Stipanuk, Martha H. (2006). Biochemical, Physiological & Molecular Aspects of Human Nutrition. W. B. Saunders Company. pp. 1043–1067. ISBN 978-0-7216-4452-3. [143] Greenwood 1997, pp. 1224–1225 [144] Kohen, Amnon; Limbach, Hans-Heinrich (2006). Isotope Effects in Chemistry and Biology (http:/ / books. google. com/ books?id=7EiIqrRBBQgC). Boca Raton, Florida: CRC Press. p. 850. ISBN 0824724496. . [145] Greenwood 1997, p. 1225 [146] Cotton 1999, p. 627 [147] Gadallah, M. A. A. (2000). "Effects of indole-3-acetic acid and zinc on the growth, osmotic potential and soluble carbon and nitrogen components of soybean plants growing under water deficit". Journal of Arid Environments 44: 451. doi: 10.1006/jare.1999.0610 (http:/ / dx. doi. org/ 10. 1006/ jare. 1999. 0610). [148] Cotton 1997, p. 628 [149] Whitney, Eleanor Noss; Rolfes, Sharon Rady (2005). Understanding Nutrition (10th ed.). Thomson Learning. pp. 447–450. ISBN 978-1-4288-1893-4. [150] NRC 2000, p. 447 [151] Hershfinkel, Michal; Silverman, William F.; Sekler, Israel (2007). "The Zinc Sensing Receptor, a Link Between Zinc and Cell Signaling". Molecular Medicine 13: 331. doi: 10.2119/2006–00038.Hershfinkel. (http:/ / dx. doi. org/ 10. 2119/ 2006â00038. Hershfinkel. ). PMID 17728842. [152] Cotton 1999, p. 629 [153] Blake, Steve (2007). Vitamins and Minerals Demystified. McGraw-Hill Professional. p. 242. ISBN 0-07-148901-0. [154] Fosmire, G. J. (1990). " Zinc toxicity (http:/ / www. ajcn. org/ cgi/ content/ abstract/ 51/ 2/ 225)". American Journal of Clinical Nutrition 51: 225. PMID 2407097. . [155] NRC 2000, p. 442 [156] Ensminger, Audrey H.; Konlande, James E. (1993). Foods & Nutrition Encyclopedia (http:/ / books. google. com/ books?id=XMA9gYIj-C4C) (2nd ed.). Boca Raton, Florida: CRC Press. pp. 2368–2369. ISBN 0849389801. . [157] " Zinc content of selected foods per common measure (http:/ / www. nal. usda. gov/ fnic/ foodcomp/ Data/ SR20/ nutrlist/ sr20w309. pdf)" (PDF). USDA National Nutrient Database for Standard Reference, Release 20. United States Department of Agriculture. . Retrieved 2007-12-06. [158] Allen, Lindsay H. (1998). " Zinc and micronutrient supplements for children (http:/ / www. ajcn. org/ cgi/ reprint/ 68/ 2/ 495S)". American Journal of Clinical Nutrition 68 (2 Suppl): 495S. PMID 9701167. . [159] Maret, W.; Sandstead H. H. (2006). "Zinc requirements and the risks and benefits of zinc supplementation". Journal of Trace Elements in Medicine and Biology 20: 3. doi: 10.1016/j.jtemb.2006.01.006 (http:/ / dx. doi. org/ 10. 1016/ j. jtemb. 2006. 01. 006). PMID 16632171. [160] Rosado, J. L. (2003). "Zinc and copper: proposed fortification levels and recommended zinc compounds". Journal of Nutrition 133: 2985S. PMID 12949397. [161] Hotz, C.; DeHaene, J.; Woodhouse, L. R.; Villalpando, S.; Rivera, J. A.; King, J. C. (2005). " Zinc absorption from zinc oxide, zinc sulfate, zinc oxide + EDTA, or sodium-zinc EDTA does not differ when added as fortificants to maize tortillas (http:/ / jn. nutrition. org/ cgi/ pmidlookup?view=long& pmid=15867288)". Journal of
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Zinc Nutrition 135: 1102. PMID 15867288. . [162] Ibs, K. H.; Rink, L.; (2003). " Zinc-altered immune function (http:/ / jn. nutrition. org/ cgi/ pmidlookup?view=long& pmid=12730441)". Journal of Nutrition 133 (5 Suppl 1): 1452S. PMID 12730441. . [163] Hambidge, M. (2003). " Biomarkers of trace mineral intake and status (http:/ / jn. nutrition. org/ cgi/ pmidlookup?view=long& pmid=12612181)". Journal of Nutrition 133 Suppl 3: 948S. PMID 12612181. . [164] WHO contributors (2007). " The impact of zinc supplementation on childhood mortality and severe morbidity (http:/ / www. who. int/ child_adolescent_health/ documents/ zinc_mortality/ en/ index. html)" (PDF). World Health Organization. . Retrieved 2009-03-01. [165] Shrimpton, R.; Gross, R.; Darnton-Hill, I.; Young, M. (2005). " Zinc deficiency: what are the most appropriate interventions? (http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?tool=pmcentrez& artid=548733)". British Medical Journal 330: 347. doi: 10.1136/bmj.330.7487.347 (http:/ / dx. doi. org/ 10. 1136/ bmj. 330. 7487. 347). PMID 15705693. [166] Eisler, Ronald (1993). " Zinc Hazard to Fish, Wildlife, and Invertebrates: A Synoptic Review (http:/ / www. pwrc. usgs. gov/ infobase/ eisler/ chr_26_zinc. pdf)" (PDF). Contaminant Hazard Reviews (Laurel, Maryland: U.S. Department of the Interior, Fish and Wildlife Service) (10). . [167] Muyssen, Brita, T. A.; De Schamphelaere, Karel A. C.; Janssen, Colin R. (2006). "Mechanisms of chronic waterborne Zn toxicity in Daphnia magna". Aquatic Toxicology 77: 393. doi: 10.1016/j.aquatox.2006.01.006 (http:/ / dx. doi. org/ 10. 1016/ j. aquatox. 2006. 01. 006). [168] Bothwell, Dawn N.; Mair, Eric A.; Cable, Benjamin B. (2003). "Chronic Ingestion of a Zinc-Based Penny". Pediatrics 111: 689. doi: 10.1542/peds.111.3.689 (http:/ / dx. doi. org/ 10. 1542/ peds. 111. 3. 689). PMID 12612262. [169] Johnson AR, Munoz A, Gottlieb JL, Jarrard DF (2007). "High dose zinc increases hospital admissions due to genitourinary complications". J. Urol. 177: 639. doi: 10.1016/j.juro.2006.09.047 (http:/ / dx. doi. org/ 10. 1016/ j. juro. 2006. 09. 047). PMID 17222649. [170] FDA says Zicam nasal products harm sense of smell (http:/ / www. latimes. com/ news/ nationworld/ nation/ la-sci-zicam17-2009jun17,0,3013664. story), Los Angeles Times, June 17, 2009 [171] Barceloux, Donald G.; Barceloux, Donald (1999). "Zinc". Clinical Toxicology 37: 279. doi: 10.1081/CLT-100102426 (http:/ / dx. doi. org/ 10. 1081/ CLT-100102426). [172] Bennett, Daniel R. M.D.; Baird, Curtis J. M.D.; Chan, Kwok-Ming; Crookes, Peter F.; Bremner, Cedric G.; Gottlieb, Michael M.; Naritoku, Wesley Y. M.D. (1997). "Zinc Toxicity Following Massive Coin Ingestion.". American Journal of Forensic Medicine & Pathology 18: 148. doi: 10.1097/00000433-199706000-00008 (http:/ / dx. doi. org/ 10. 1097/ 00000433-199706000-00008). [173] Fernbach, S. K.; Tucker G. F. (1986). " Coin ingestion: unusual appearance of the penny in a child (http:/ / radiology. rsnajnls. org/ cgi/ content/ abstract/ 158/ 2/ 512)". Radiology 158: 512. PMID 3941880. . [174] Stowe, C. M.; Nelson, R.; Werdin, R.; et al. (1978). "Zinc phosphide poisoning in dogs". Journal of the American Veterinary Medical Association 173: 270. PMID 689968. [175] Reece, R. L.; Dickson, D. B.; Burrowes, P. J. (1986). "Zinc toxicity (new wire disease) in aviary birds". Australian Veterinary Journal 63: 199. doi: 10.1111/j.1751-0813.1986.tb02979.x (http:/ / dx. doi. org/ 10. 1111/ j. 1751-0813. 1986. tb02979. x).
Bibliography • Chambers, William and Robert (1901). Chambers's Encyclopaedia: A Dictionary of Universal Knowledge (http:/ / books. google. com/ books?id=Rz8oAAAAYAAJ& printsec=toc& client=firefox-a& source=gbs_summary_r& cad=0) (Revised ed.). London and Edinburgh: J. B. Lippincott Company. http:/ / books. google. com/ books?id=Rz8oAAAAYAAJ& printsec=toc& client=firefox-a& source=gbs_summary_r& cad=0. • Cotton, F. Albert; Wilkinson, Geoffrey; Murillo, Carlos A.; Bochmann, Manfred (1999). Advanced Inorganic Chemistry (6th ed.). New York: John Wiley & Sons, Inc.. ISBN 0-471-199957-5. • CRC contributors (2006). David R. Lide. ed. Handbook of Chemistry and Physics (87th ed.). Boca Raton, Florida: CRC Press, Taylor & Francis Group. ISBN 0-8493-0487-3. • Emsley, John (2001). " Zinc (http:/ / books. google. com/ books?id=j-Xu07p3cKwC)". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. pp. 499–505. ISBN 0-19-850340-7. http:/ / books. google. com/
24
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25 books?id=j-Xu07p3cKwC.
• Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. ISBN 0750633654. • Heiserman, David L. (1992). " Element 30: Zinc (http:/ / books. google. com/ books?id=24l-Cpal9oIC& pgis=1)". Exploring Chemical Elements and their Compounds. New York: TAB Books. ISBN 0-8306-3018-X. http:/ / books. google. com/ books?id=24l-Cpal9oIC& pgis=1. • Lehto, R. S. (1968). "Zinc". in Clifford A. Hampel. The Encyclopedia of the Chemical Elements. New York: Reinhold Book Corporation. pp. 822–830. LCCN 68-29938. ISBN 0-442-15598-0. • United States National Research Council, Institute of Medicine. (2000). Dietary Reference Intakes for Vitamin A, Vitamin K, Arsenic, Boron, Chromium, Copper, Iodine, Iron, Manganese, Molybdenum, Nickel, Silicon, Vanadium, and Zinc (http:/ / www. nap. edu/ catalog. php?record_id=10026). National Academies Press. pp. 442–455. http:/ / www. nap. edu/ catalog. php?record_id=10026. • Stwertka, Albert (1998). "Zinc". Guide to the Elements (Revised ed.). Oxford University Press. ISBN 0195080831. • Weeks, Mary Elvira (1933). "III. Some Eighteenth-Century Metals". The Discovery of the Elements. Easton, PA: Journal of Chemical Education. ISBN 0766138720.
External links • WebElements.com – Zinc (http:/ / www. webelements. com/ webelements/ elements/ text/ Zn/ index. html) • History & Etymology of Zinc (http:/ / www. vanderkrogt. net/ elements/ elem/ zn. html) • Statistics and Information from the U.S. Geological Survey (http:/ / minerals. er. usgs. gov/ minerals/ pubs/ commodity/ zinc/ ) • Reducing Agents > Zinc (http:/ / www. organic-chemistry. org/ chemicals/ reductions/ zinc-zn. shtm) • American Zinc Association (http:/ / www. zinc. org) Information about the uses and properties of zinc. • Outline safety data for zinc (http:/ / ptcl. chem. ox. ac. uk/ MSDS/ ZI/ zinc. html)
Article Sources and Contributors
Article Sources and Contributors Zinc Source: http://en.wikipedia.org/w/index.php?oldid=307717156 Contributors: -jmac-, 194.200.130.xxx, 21655, 2D, 56, Abhisht, Acroterion, Adam Johnston, Adashiel, Adrian, Ahoerstemeier, Aitias, Akadruid, Alansohn, AlcheMister, Alchemist01010101, Ale jrb, AlexR, AlexiusHoratius, Alphachimp, AndonicO, Andres, Andrewrost3241981, Animum, Anlace, Annabanana308, Antandrus, Anwar saadat, Aphaia, Arcadian, Archimerged, ArglebargleIV, Ashley beauchamp, AssegaiAli, Auric, AutumnalMelody, Awils1, AxelBoldt, Axiosaurus, Axl, Bantman, Bart133, Basicdesign, Beano, BeastmasterGeneral, Beetstra, Bellatrix Kerrigan, BenFrantzDale, BerserkerBen, Betterusername, Bettia, Bhadani, Bharatveer, Bhendrickson, Blanchardb, Blehfu, BlueEarth, Bobo192, Bogdana, Bogdangiusca, Bomac, Bongwarrior, Bookwyrm404, Borislav.dopudja, Bork, BostonMA, Boyd Steere, Bradtcordeiro, Brady
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27
Gallium
1
Gallium zinc ← gallium → germaniumAl ↑ Ga ↓ In
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
31Ga Periodic table
Appearance silvery white
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
Gallium
2 at T/K
1310
1448
1620
1838
2125
2518
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1979.3 kJ·mol−1 3rd: 2963 kJ·mol−1Atomic radiusCovalent radiusVan der Waals radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityThermal expansionSpeed of sound (thin rod) Young's modulusPoisson ratioMohs hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of gallium iso
N.A.
half-life
69
60.11%
69
71
39.89%
71
Ga Ga
DM
DE (MeV)
DP
Ga is stable with 38 neutron Ga is stable with 40 neutron
gallium, Ga, 31 poor metal13, 4, d69.723(1) g·mol−1 [Ar] 3d10 4s2 4p1 2, 8, 18, 3 (Image) solid 5.91 g·cm−3 6.095 g·cm−3 302.9146 K,29.7646 °C,85.5763 °F 2477 K,2204 °C,3999 °F 5.59 kJ·mol−1 254 kJ·mol−1 (25 °C) 25.86 J·mol−1·K−13, 2, 1 (amphoteric oxide) 1.81 (Pauling scale) 1st: 578.8 kJ·mol−1 135 pm 122±3 pm 187 pm orthorhombic diamagnetic (20 °C) 270Ω·m (300 K) 40.6 W·m−1·K−1 (25 °C) 1.2 µm·m−1·K−1 (20 °C) 2740 m/s 9.8 GPa 0.47 1.5 60 MPa 7440-55-3 Gallium (pronounced /ˈɡæliəm/) is a chemical element that has the symbol Ga and atomic number 31. Elemental gallium does not occur in nature, but as the Ga (III) salt, in trace amounts in bauxite and zinc ores. A soft silvery metallic poor metal, elemental gallium is a brittle solid at low temperatures. As it liquefies slightly above room temperature, it will melt in the hand. Its melting point is used as a temperature reference point, and from its discovery in 1875 to the semiconductor era, its primary uses were in high-temperature thermometric applications and in preparation of metal alloys with unusual properties of stability, or ease of melting; some being liquid at room temperature or below. The alloy Galinstan (68.5% Ga, 21.5% In, 10% Sn) has a melting point of about -19 °C (-2.2 °F). In semiconductors, an important application is in the compounds gallium arsenide and gallium nitride, used most notably in light-emitting diodes (LEDs). Semiconductor use is now almost the entire (> 95%) world market for gallium, but new uses in alloys and fuel cells continue to be discovered. Gallium is not known to be essential in biology, but because of the biological handling of gallium’s primary ionic salt Ga(III) as though it were iron(III), gallium ion localizes to and interacts with many processes in the body in which iron(III) is manipulated. As these processes include inflammation, which is present as a marker for many disease states, several gallium salts are used, or are in development, as both pharmaceuticals and radiopharmaceuticals in medicine.
Notable characteristics Elemental gallium is not found in nature, but it is easily obtained by smelting. Very pure gallium metal has a brilliant silvery color and its solid metal fractures conchoidally like glass. Gallium metal expands by 3.1 percent when it solidifies, and therefore storage in either glass or metal containers is avoided, due to the possibility of container rupture with freezing. Gallium shares the higher-density liquid state with only a few materials like silicon, germanium, bismuth, antimony and water.
Gallium
3
Gallium also attacks most other metals by diffusing into their metal lattice. Gallium for example diffuses into the grain boundaries of Al/Zn alloys[1] or steel[2] , making them very brittle. Also, gallium metal easily alloys with many metals, and was used in small quantities in the core of the first atomic bomb to help stabilize the plutonium crystal structure.[3] The melting point of 302.9146 K (29.7646°C, 85.5763°F) is near room temperature. Gallium's melting point (mp) is one of the formal temperature reference points in the International Temperature Scale of 1990 (ITS-90) established by BIPM.[4] [5] [6] The triple point of gallium of 302.9166 K (29.7666°C, 85.5799°F), is being used by NIST in preference to gallium's melting point.[7] Gallium is a metal that will melt in one's hand. This metal has a strong tendency to supercool below its melting point/freezing point. Seeding with a crystal helps to initiate freezing. Gallium is one of the metals (with caesium, rubidium, francium and mercury) which are liquid at or near normal room temperature, and can therefore be used in metal-in-glass high-temperature thermometers. It is also notable for having one of the largest liquid ranges for a metal, and (unlike mercury) for having a low vapor pressure at high temperatures. Unlike mercury, liquid gallium metal wets glass and skin, making it mechanically more difficult to handle (even though it is substantially less toxic and requires far fewer precautions). For this reason as well as the metal contamination problem and freezing-expansion problems noted above, samples of gallium metal are usually supplied in polyethylene packets within other containers. Gallium does not crystallize in any of the simple crystal structures. The stable phase under normal conditions is orthorhombic with 8 atoms in the conventional unit cell. Each atom has only one nearest neighbor (at a distance of 244 pm) and six other neighbors within additional 39 pm. Many stable and metastable phases are found as function of temperature and pressure. The bonding between the nearest neighbors is found to be of covalent character, hence Ga2 dimers are seen as the fundamental building blocks of the crystal. This explains the drop of the melting point compared to its neighbour elements aluminium and indium. The compound with arsenic, gallium arsenide is a semiconductor commonly used in light-emitting diodes. Crystallization of gallium from the melt
High-purity gallium is dissolved slowly by mineral acids. Gallium has no known biological role, although it has been observed to stimulate metabolism.[8]
History Gallium (the Latin Gallia means "Gaul," essentially modern France) was discovered spectroscopically by Lecoq de Boisbaudran in 1875 by its characteristic spectrum (two violet lines) in an examination of a zinc blende from the Pyrenees.[9] Before its discovery, most of its properties had been predicted and described by Dmitri Mendeleev (who had called the hypothetical element "eka-aluminium" on the basis of its position in his periodic table). Later, in 1875, Boisbaudran obtained the free metal by electrolysis of its hydroxide in potassium hydroxide solution. He named the element "gallia" after his native land of
Gallium France. It was later claimed that, in one of those multilingual puns so beloved of men of science in the early 19th century, he had also named gallium after himself, as his name, "Le coq," is the French for "the rooster," and the Latin for "rooster" is "gallus"; however, in an 1877 article Le coq denied this supposition. (The supposition was also noted in Building Blocks of the Universe, a book on the elements by Isaac Asimov.)
Occurrence Gallium does not exist in free form in nature, and the few high-gallium minerals such as gallite (CuGaS2) are too rare to serve as a primary source of the element or its compounds. Its abundance in the Earth's crust is approximately 16.9 ppm.[10] Gallium is found and extracted as a trace component in bauxite and to a small extent from sphalerite. The amount extracted from coal, diaspore and germanite in which gallium is also present is negligible. The United States Geological Survey (USGS) estimates gallium reserves to exceed 1 million tonnes, based on 50 ppm by weight concentration in known reserves of bauxite and zinc ores.[11] [12] Some flue dusts from burning coal have been shown to contain small quantities of gallium, typically less than 1% by weight.[13] [14] [15] [16]
Production The only two economic sources for gallium are as byproduct of aluminium and zinc production, while the sphalerite for zinc production is the minor source. Most gallium is extracted from the crude aluminium hydroxide solution of the Bayer process for producing alumina and aluminium. A mercury cell electrolysis and hydrolysis of the amalgam with sodium hydroxide leads to sodium gallate. Electrolysis then gives gallium metal. For semiconductor use, further purification is carried out using zone melting, or else single crystal extraction from a melt (Czochralski process). Purities of 99.9999% are routinely achieved and commercially widely available.[17] An exact number for the world wide production is not available, but it is estimated that in 2007 the production of gallium was 184 tonnes with less than 100 tonnes from mining and the rest from scrap recycling. [11]
Applications Semiconductors Gallium arsenide (GaAs) and gallium nitride (GaN) used in electronic components represented about 98% of the gallium consumption in the United States.[11] World wide gallium arsenide makes up 95% of the annual global gallium consumption.[17] The semiconductor applications are the main reason for the low-cost commercial availability of the extremely high-purity (99.9999+%) metal: As a component of the semiconductor gallium arsenide, the most common Gallium based blue LEDs application for gallium is optoelectronic devices (mostly laser diodes and light-emitting diodes.) Smaller amounts of gallium arsenide are use for the manufacture of ultra-high speed logic chips and MESFETs for low-noise microwave preamplifiers.
4
Gallium Gallium is used as a dopant for the production of solid-state devices such as transistors. However, worldwide the actual quantity used for this purpose is minute, since dopant levels are usually of the order of a few parts per million. Multijunction photovoltaic cell is used for special application, first developed and deployed for satellite power applications, are made by molecular beam epitaxy or Metalorganic vapour phase epitaxy of thin films of gallium arsenide, indium gallium phosphide or indium gallium arsenide.The Mars Exploration Rovers and several satellites use triple junction gallium arsenide on germanium cells.[18] Gallium is the rarest component of new photovoltaic compounds (such as copper indium gallium selenium sulfide or Cu(In,Ga)(Se,S)2) for use in solar panels as a more efficient alternative to crystalline silicon.[19]
Wetting and alloy improvement • Because gallium wets glass or porcelain, gallium can be used to create brilliant mirrors. • Gallium readily alloys with most metals, and has been used as a component in low-melting alloys. The plutonium used in nuclear weapon pits is machined by alloying with gallium to stabilize the allotropes of plutonium.[20] • Gallium added in quantities up to 2% in common solders can aid wetting and flow characteristics.
Galinstan and other liquid alloys A nearly eutectic alloy of gallium, indium, and tin is a room temperature liquid which is widely available in medical thermometers, replacing problematic mercury. This alloy, with the trade-name Galinstan (with the "-stan" referring to the tin), has a low freezing point of -19 °C (-2.2°F).[21] It has been suggested that this family of alloys could also be used to cool computer chips in place of water.[22] Much research is being devoted to gallium alloys as substitutes for mercury dental amalgams, but these compounds have yet to see wide acceptance.
Energy storage Aluminium is reactive enough to reduce water to hydrogen, being oxidized to aluminium oxide. However, the aluminium oxide forms a protective coat which prevents further reaction. Galinstan has been applied to activate aluminum (removing the oxide coat), so that aluminum can react with water, generating hydrogen and steam in a reaction being considered as a helpful step in a hydrogen economy[23] . A number of other gallium-alluminum alloys are also usable for the purpose of essentially acting as chemical energy store to generate hydrogen from water, on-site. After reaction with water, resmelting the resultant aluminium oxide and gallium mixture to metallic aluminium and gallium might be reforming back into electrodes, with energy input. [23] [24] The thermodynamic efficiency of the aluminium smelting process is estimated as [25] 50%. Therefore, at most no more than half the energy that goes into smelting the aluminium could be recovered by a hydrogen fuel cell.
5
Gallium
Biomedical applications As gallium (III) salts • Gallium nitrate (see Ganite) has been used as an intravenous pharmaceutical to treat hypercalcemia associated with tumor metastatis to bones. Gallium is thought to interfere with osteoclast function. It may be effective when other treatments for maligancy-associated hypercalcemia are not. [26] • Gallium maltolate is in clinical and preclinical trials as a potential treatment for cancer, infectious disease, and inflammatory disease.[27] • Research is being conducted to determine whether gallium can be used to fight bacterial infections in people with cystic fibrosis. Gallium is similar in size to iron, an essential nutrient for respiration. When gallium is mistakenly picked up by bacteria such as Pseudomonas, the bacteria's ability to respire is interfered with and the bacteria die. The mechanism behind this is that iron is redox active, which allows for the transfer of electrons during respiration, but gallium is redox inactive.[28] [29] As radiogallium salts Gallium-67 salts such as gallium citrate and gallium nitrate are used as radiopharmaceutical agents in a nuclear medicine imaging procedure commonly referred to as a gallium scan. The form or salt of gallium is not important, since it is the free dissolved gallium ion Ga3+ which is the active radiotracer. For these applications, the radioactive isotope 67Ga is used. The body handles Ga3+ in many ways as though it were iron, and thus it is bound (and concentrates) in areas of inflammation, such as infection, and also areas of rapid cell division. This allows such sites to be imaged by nuclear scan techniques. This use has largely been replaced by fluorodeoxyglucose (FDG) for positron emission tomography, "PET" scan and indium-111 labelled leukocyte scans. However, the localization of gallium in the body has some properties which make it unique in some circumstances from competing modalities using other radioisotopes. Gallium-68, a positron emitter with a half life of 68 min., is now used as a diagnostic radionuclide in CT-PET when linked to pharmaceutical preparations such as DOTATOC, a somatostatin analogue used for neuroendocrine tumors investigation, and DOTATATE, a newer one, used for neuroendocrine metastasis and lung neuroendocrine cancer, such as certain types of microcytoma. Galium-68's preparation as a pharmaceutical is chemical and the radionuclide is extracted by elution from germanium-68, a synthetic radioisotope of germanium, in gallium-68 generators. These generators function similarly to technetium-99m generators, in both cases using a process similar to thin layer chromatography. The stationary phase is alumina, TiO2 or SnO2, onto which germanium-68 is adsorbed. The mobile phase is a solvent able to elute (wash out) decayed germanium-68, after it has decayed to gallium-68 (III). Currently Ga-68 is easily eluted with a few mL of 1 M or 0.1M hydrochloric acid from tin-oxide or titanium-oxide based generators respectively, within 1 to 2 minutes. However, there remains more than an hour of pharmaceutical preparation to attach the gallium-68 (III) to the tracer DOTATOC or DOTATATE, so that the total preparation time is typically longer than the Ga-68 isotope half life. This fact requires that these radiopharmaceuticals be made on-site in most cases. The on-site generator is required to minimize the time losses. The generator is easily storable for almost a year.
6
Gallium
7
Other uses • Magnesium gallate containing impurities (such as Mn2+), is beginning to be used in ultraviolet-activated phosphor powder. • Neutrino detection. Possibly the largest amount of pure gallium ever collected in a single spot is the Gallium-Germanium Neutrino Telescope used by the SAGE experiment at the Baksan Neutrino Observatory in Russia. This detector contains 55-57 tonnes of liquid gallium.[30] Another experiment was the GALLEX neutrino detector operated in the early 1990s in an Italian mountain tunnel. The detector contained 12.2 tons of watered gallium-71. Solar neutrinos caused a few atoms of Ga-71 to become radioactive Ge-71, which were detected. The solar neutrino flux deduced was found to have a deficit of 40% from theory. This was not explained until better solar neutrino detectors and theories were constructed (see SNO).[31] • As a liquid metal ion source for a focused ion beam.
Precautions While not considered toxic, the data about gallium are inconclusive. Some sources suggest that it may cause dermatitis from prolonged exposure; other tests have not caused a positive reaction. Like most metals, finely divided gallium loses its luster and powdered gallium appears gray. Thus, when gallium is handled with bare hands, the extremely fine dispersion of liquid gallium droplets, which results from wetting skin with the metal, may appear as a gray skin stain.
See also • Gallium compounds
External links • • • • • • •
Webelements: detailed information on gallium [32] WebElements.com – textbook information on gallium [33] Material safety data sheet at acialloys.com [34] www.lenntech.com – textbook information regarding gallium [35] environmental effects of gallium [36] Price development of gallium 1959-1998 [37] Technology produces hydrogen by adding water to an alloy of aluminum and gallium
[38]
• pure Gallium crystals ~99,9999% picture in the element collection from Heinrich Pniok [39]
Gallium
References [1] W. L. Tsai, Y. Hwu, C. H. Chen, L. W. Chang, J. H. Je, H. M. Lin, G. Margaritondo (2003). "Grain boundary imaging, gallium diffusion and the fracture behavior of Al–Zn Alloy – An in situ study". Nuclear Instruments and Methods in Physics Research Section B 199: 457. doi: 10.1016/S0168-583X(02)01533-1 (http:/ / dx. doi. org/ 10. 1016/ S0168-583X(02)01533-1). [2] Vigilante, G. N., Trolano, E., Mossey, C. (June 1999). " Liquid Metal Embrittlement of ASTM A723 Gun Steel by Indium and Gallium (http:/ / stinet. dtic. mil/ oai/ oai?& verb=getRecord& metadataPrefix=html& identifier=ADA365497)". Defense Technical Information Center. . Retrieved 2009-07-07. [3] Sublette,Cary (2001-09-09). " Section 6.2.2.1 (http:/ / nuclearweaponarchive. org/ Nwfaq/ Nfaq6. html#nfaq6. 2)". Nuclear Weapons FAQ. . Retrieved 2008-01-24. [4] Preston=Thomas, H. (1990). " The International Temperature Scale of 1990 (ITS-90) (http:/ / www. bipm. org/ utils/ common/ pdf/ its-90/ ITS-90_metrologia. pdf)". Metrologia 27: 3–10. doi: 10.1088/0026-1394/27/1/002 (http:/ / dx. doi. org/ 10. 1088/ 0026-1394/ 27/ 1/ 002). . [5] " ITS-90 documents at Bureau International de Poids et Mesures (http:/ / www. bipm. org/ en/ publications/ its-90. html)". . [6] Magnum, B.W.; Furukawa, G.T. (August 1990). " Guidelines for Realizing the International Temperature Scale of 1990 (ITS-90) (http:/ / www. cstl. nist. gov/ div836/ 836. 05/ papers/ magnum90ITS90guide. pdf)". National Institute of Standards and Technology. NIST TN 1265. . [7] Strouse, Gregory F. (1999). " NIST realization of the gallium triple point (http:/ / www. cstl. nist. gov/ div836/ 836. 05/ papers/ Strouse99GaTP. pdf)". National Institute of Standards and Technology. . Retrieved 2009-07-07. [8] Mark Winter. " Scholar Edition: gallium: Biological information (http:/ / www. webelements. com/ webelements/ scholar/ elements/ gallium/ biological. html)". The University of Sheffield and WebElements Ltd, UK. . [9] de Boisbaudran, Lecoq. " Caractères chimiques et spectroscopiques d'un nouveau métal, le gallium, découvert dans une blende de la mine de Pierrefitte, vallée d'Argelès (Pyrénées) (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k3038w/ f490. table)". Comptes rendus 81: 493. . Retrieved 2008-09-23. [10] Burton, J. D.; Culkin, F.; Riley, J. P. (2007). "The abundances of gallium and germanium in terrestrial materials". Geochimica et Cosmochimica Acta 16: 151. doi: 10.1016/0016-7037(59)90052-3 (http:/ / dx. doi. org/ 10. 1016/ 0016-7037(59)90052-3). [11] Kramer, Deborah A.. " Mineral Commodity Summary 2006: Gallium (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ gallium/ mcs-2008-galli. pdf)". United States Geological Survey. . Retrieved 2008-11-20. [12] Kramer, Deborah A.. " Mineral Yearbook 2006: Gallium (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ gallium/ myb1-2006-galli. pdf)". United States Geological Survey. . Retrieved 2008-11-20. [13] Shan Xiao-quan, Wang Wen and Wen Bei (1992). "Determination of gallium in coal and coal fly ash by electrothermal atomic absorption spectrometry using slurry sampling and nickel chemical modification". Journal of Analytical Atomic Spectrometry 7: 761. doi: 10.1039/JA9920700761 (http:/ / dx. doi. org/ 10. 1039/ JA9920700761). [14] " Gallium in West Virginia Coals (http:/ / www. wvgs. wvnet. edu/ www/ datastat/ te/ GaHome. htm)". West Virginia Geological and Economic Survey. 2002-03-02. . [15] O. Font, X. Querol, R. Juan, R. Casado, C. R. Ruiz, A. Lopez-Soler, P. Coca and F. G. Pena (2007). "Recovery of gallium and vanadium from gasification fly ash". Journal of Hazardous Materials 139: 413. doi: 10.1016/j.jhazmat.2006.02.041 (http:/ / dx. doi. org/ 10. 1016/ j. jhazmat. 2006. 02. 041). [16] A. J. W. Headlee and Richard G. Hunter (1953). "Elements in Coal Ash and Their Industrial Significance". Industrial and Engineering Chemistry 45: 548. doi: 10.1021/ie50519a028 (http:/ / dx. doi. org/ 10. 1021/ ie50519a028). [17] Moskalyk, R. R. (2003). "Gallium: the backbone of the electronics industry". Minerals Engineering 16: 921. doi: 10.1016/j.mineng.2003.08.003 (http:/ / dx. doi. org/ 10. 1016/ j. mineng. 2003. 08. 003). [18] Crisp, D.; Pathare, A.; Ewell, R. C. (2004). "The performance of gallium arsenide/germanium solar cells at the Martian surface". Progress in Photovoltaics Research and Applications 54: 83. doi: 10.1016/S0094-5765(02)00287-4 (http:/ / dx. doi. org/ 10. 1016/ S0094-5765(02)00287-4). [19] Alberts, V.; Titus J.; Birkmire R. W. (2003). "Material and device properties of single-phase Cu(In,Ga)(Se,S)2 alloys prepared by selenization/sulfurization of metallic alloys". Thin Solid Films 451-452: 207. doi: 10.1016/j.tsf.2003.10.092 (http:/ / dx. doi. org/ 10. 1016/ j. tsf. 2003. 10. 092). [20] Besmann, Theodore M. (2005). "Thermochemical Behavior of Gallium in Weapons-Material-Derived Mixed-Oxide Light Water Reactor (LWR) Fuel". Journal of the American Ceramic Society 81: 3071. doi: 10.1111/j.1151-2916.1998.tb02740.x (http:/ / dx. doi. org/ 10. 1111/ j. 1151-2916. 1998. tb02740. x). [21] Surmann, P; Zeyat, H (Nov 2005). "Voltammetric analysis using a self-renewable non-mercury electrode.". Analytical and bioanalytical chemistry 383 (6): 1009–13. doi: 10.1007/s00216-005-0069-7 (http:/ / dx. doi. org/
8
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9
10. 1007/ s00216-005-0069-7). ISSN 1618-2642 (http:/ / worldcat. org/ issn/ 1618-2642). PMID 16228199. [22] Knight, Will (2005-05-05). " Hot chips chilled with liquid metal (http:/ / www. newscientist. com/ article. ns?id=dn7348)". . Retrieved 2008-11-20. [23] Purdue University (2007-04-10). " Purdue Energy Center symposium to pave the road to a hydrogen economy (http:/ / www. purdue. edu/ uns/ x/ 2007a/ 070410Gorehydrogen. html)". Press release. . [24] " New process generates hydrogen from aluminum alloy to run engines, fuel cells (http:/ / www. physorg. com/ news98556080. html)". PhysOrg.com. 2007-05-16. . [25] Das, Subodh K. (2004). "Energy implications of the changing world of aluminum metal supply". JOM 56: 14. doi: 10.1007/s11837-004-0175-6 (http:/ / dx. doi. org/ 10. 1007/ s11837-004-0175-6). [26] " gallium nitrate (http:/ / www. cancer. org/ docroot/ CDG/ content/ CDG_gallium_nitrate. asp)". . Retrieved 2009-07-07. [27] L. R. Bernstein, T. Tanner, C. Godfrey, B. Noll (2000). "Chemistry and pharmacokinetics of gallium maltolate, a compound with high oral gallium bioavailability". Metal Based Drugs 7: 33. doi: 10.1155/MBD.2000.33 (http:/ / dx. doi. org/ 10. 1155/ MBD. 2000. 33). [28] " A Trojan-horse strategy selected to fight bacteria (http:/ / www. infoniac. com/ health-fitness/ trojan-gallium. html)". 2007-03-16. . Retrieved 2008-11-20. [29] Smith, Michael (2007-03-16). " Gallium May Have Antibiotic-Like Properties (http:/ / www. medpagetoday. com/ InfectiousDisease/ GeneralInfectiousDisease/ tb/ 5266)". MedPage Today. . Retrieved 2008-11-20. [30] " Russian American Gallium Experiment (http:/ / ewi. npl. washington. edu/ sage/ )". 2001-19-10. . Retrieved 2009-6-24. [31] " Neutrino Detectors Experiments: GALLEX (http:/ / wwwlapp. in2p3. fr/ neutrinos/ anexp. html#gallex)". 1999-06-26. . Retrieved 2008-11-20. [32] [33] [34] [35] [36] [37] [38] [39]
http:/ / www. webelements. com/ webelements/ elements/ text/ Ga/ key. html http:/ / www. webelements. com/ webelements/ elements/ text/ Ga/ index. html http:/ / www. acialloys. com/ msds/ ga. html http:/ / www. lenntech. com/ Periodic-chart-elements/ Ga-en. htm http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ gallium/ index. html http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ gallium/ 460798. pdf http:/ / www. physorg. com/ news107446364. html http:/ / www. pse-mendelejew. de/ bilder/ ga. jpg
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10
Germanium
1
Germanium 32
gallium ← germanium → arsenic
Si ↑
Ge ↓
Sn Periodic Table - Extended Periodic Table
General Name, Symbol, Number
germanium, Ge, 32
Element category
metalloids
Group, Period, Block
14, 4, p
Appearance
grayish white
Standard atomic weight
72.64(1) g·mol
Electron configuration
[Ar] 3d
−1
10
Electrons per shell
2
2
4s 4p
2, 8, 18, 4 Physical properties
Phase
solid
Density (near r.t.)
5.323 g·cm−3
Liquid density at m.p.
5.60 g·cm−3
Melting point
1211.40 K (938.25 °C, 1720.85 °F)
Boiling point
3106 K (2833 °C, 5131 °F)
Heat of fusion
36.94 kJ·mol−1
Heat of vaporization
334 kJ·mol−1
Specific heat capacity
(25 °C) 23.222 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1644
1814
2023
2287
2633
3104
Atomic properties
Germanium
2
Crystal structure
Diamond cubic
Oxidation states
4, 3, 2, 1, 0, -1, -2, -3, -4 (amphoteric oxide)
Electronegativity
2.01 (Pauling scale)
Ionization energies (more)
1st: 762 kJ·mol−1 2nd: 1537.5 kJ·mol−1 3rd: 3302.1 kJ·mol−1
Atomic radius
122 pm
Covalent radius
122 pm Miscellaneous [1]
Magnetic ordering
Diamagnetic
Electrical resistivity
(20 °C) 1 Ω·m
Thermal conductivity
(300 K) 60.2 W·m
Thermal expansion
(25 °C) 6.0 µm·m
Speed of sound (thin rod)
(20 °C) 5400 m/s
Young's modulus
103
Shear modulus
41
Bulk modulus
75
Poisson ratio
0.26
−1
−1
[2]
[2] [2]
−1
·K
·K
−1
GPa
GPa GPa
[2]
Mohs hardness
6.0
Vickers hardness
MPa
Brinell hardness
MPa
CAS registry number
7440-56-4 Most-stable isotopes
Main article: Isotopes of germanium iso 68
Ge
70
Ge
71
Ge
NA syn 21.23% syn
half-life 270.8 d
11.26 d 72
73
7.73%
73
74
35.94%
74
76
7.44%
1.78×1021 y
Ge Ge
DP
-
68
-
71
-
76
Ga
Ge is stable with 38 neutron
27.66%
Ge
ε
DE (MeV)
70
72
Ge
DM
ε
Ga
Ge is stable with 40 neutron Ge is stable with 41 neutron Ge is stable with 42 neutron β−β− References
Se
Germanium Germanium (pronounced /dʒərˈmeɪniəm/) is a chemical element with the symbol Ge and atomic number 32. It is a lustrous, hard, grayish-white metalloid in the carbon group, chemically similar to its group neighbors tin and silicon. Germanium has five naturally occurring isotopes ranging in atomic mass number from 70 to 76. It forms a large number of organometallic compounds, including tetraethylgermane and isobutylgermane. Because few minerals contain it in large concentration, germanium was discovered comparatively late despite the fact that it is relatively abundant in the Earth's crust. In 1869, Dmitri Mendeleev predicted its existence and some of its properties based on its position on his periodic table and called the element ekasilicon. Nearly two decades later, in 1886, Clemens Winkler found it in the mineral argyrodite. Winkler found that experimental observations agreed with Mendeleev's predictions and named the element after his country, Germany. Germanium is an important semiconductor material used in transistors and various other electronic devices. Its major end uses are fiber-optic systems and infrared optics, but it is also used for polymerization catalysts, in electronics and in solar electric applications. Germanium is mined primarily from sphalerite, though it is also recovered from silver, lead, and copper ores. Some germanium compounds, such as germanium chloride and germane, can irritate the eyes, skin, lungs, and throat.
History In his report on The Periodic Law of the Chemical Elements, in 1869, Dmitri Mendeleev predicted the existence of several unknown elements, including one filling a gap in the carbon family, between silicon and tin.[3] Because of its position in the table, he called it ekasilicon (Es) and assigned it an atomic weight of 72. In mid-1885, in a mine near Freiberg, Saxony, a new mineral was found. It was named argyrodite, because of its high silver content.[a] Clemens Winkler examined this new mineral and was able to isolate an element similar to antimony in 1886.[4] Before he published his results on the new element Winkler intended to name the element neptunium, as the actual discovery of planet Neptune in 1846 had been preceded by Dmitri Mendeleev mathematical prediction of its existence.[b] However, the name neptunium had already been given to an element (though not the element that today bears the name neptunium, discovered in 1940),[c] and instead, Winkler named the new metal germanium (from the Latin Germania for Germany) in honor of his fatherland.[4] Because the element showed similarities with the elements arsenic and antimony, its place in the periodic table was under discussion, but the similarities between Mendeleev's ekasilicon and germanium
3
Germanium
4
confirmed its place.[4] [5] With further material from 500 kg of ore from the mines in Saxony, Winkler confirmed the chemical properties of the new element in 1887.[6] [7] [8] He also determined an atomic weight of 72.32 by analyzing pure germanium tetrachloride (GeCl4), while Lecoq de Boisbaudran deduced 72.3 by a comparison of the lines in the spark spectrum of the element.[9] Winkler was able to prepare several new compounds of germanium, including the fluorides, chlorides, sulfides, germanium dioxide, and tetraethylgermane (Ge(C2H5)4), the first organogermane.[7] The physical data from these compounds—which corresponded with Mendeleev's predictions—made the discovery an important confirmation of Mendeleev's idea of element periodicity. Here is a comparison between the prediction and Winkler's data:[7] Property
Clemens Winkler
Ekasilicon
Germanium
atomic mass
72
72.59
density (g/cm3)
5.5
5.35
melting point (°C)
high
947
color
gray
gray
oxide type
refractory dioxide
refractory dioxide
oxide density (g/cm3)
4.7
4.7
oxide activity
feebly basic
feebly basic
chloride boiling point (°C)
under 100
86 (GeCl4)
chloride density (g/cm3)
1.9
1.9
Until the late 1930s, germanium was believed to be a poorly conducting metal.[10] It did not become economically significant until after 1945, when its properties as a semiconductor were recognized as being valuable in electronics. It was only during World War II, in 1941, that germanium diodes began to supplant vacuum tubes in electronic devices.[11] [12] Its first major use was the point contact Schottky diodes for radar reception during WWII.[10] The first silicon-germanium alloys were obtained in 1955.[13] Before 1945, only a few hundred kilograms of the element were produced each year, but by the end of the 1950s, annual worldwide production had reached 40 metric tons.[14] The development of the germanium transistor in 1948[15] opened the door to countless applications of solid state electronics.[16] From 1950 through the early 1970s, this area provided an increasing market for germanium, but then high purity silicon began replacing germanium in transistors, diodes, and rectifiers.[17] Silicon has superior electrical properties, but requires much higher purity—a purity which could not be commercially achieved in the early days.[18]
Germanium Meanwhile, demand for germanium in fiber optics communication networks, infrared night vision systems, and polymerization catalysts increased dramatically.[14] These end uses represented 85% of worldwide germanium consumption in 2000.[17] The U.S. government even designated germanium as a strategic and critical material, calling for a 146 ton (132 t) supply in the national defense stockpile in 1987.[14] Germanium differs from silicon in that the supply of silicon is only limited by production capacity, while that for germanium is limited by the availability of exploitable sources. As a result, while silicon could be bought in 1998 for less than $10 per kg,[14] the price of 1 kg of germanium was then almost $800.[14]
Characteristics Under standard conditions germanium is a brittle, silvery-white, semi-metallic element.[19] This form constitutes an allotrope technically known as α-germanium, which has a metallic luster and a diamond cubic crystal structure, the same as diamond.[17] At pressures above 120 kbar, a different allotrope known as β-germanium forms, which has the same structure as β-tin.[20] Along with silicon, gallium, bismuth, antimony, and water, it is one of the few substances that expands as it solidifies (i.e. freezes) from its molten state.[20] Germanium is a semiconductor. Zone refining techniques have led to the production of crystalline germanium for semiconductors that has an impurity of only one part in 1010,[21] making it one of the purest materials ever obtained.[22] The first metallic material discovered (in 2005) to become a superconductor in the presence of an extremely strong electromagnetic field was an alloy of germanium with uranium and rhodium.[23] Pure germanium is known to spontaneously extrude very long screw dislocations, referred to as germanium whiskers. The growth of these whiskers is one of the primary reasons for the failure of older diodes and transistors made from germanium; depending on what they eventually touch, they may lead to an electrical short.[24]
Chemistry Elemental germanium oxidizes slowly to GeO2 at 250 °C.[25] Germanium is insoluble in dilute acids and alkalis but dissolves slowly in concentrated sulfuric acid and reacts violently with molten alkalis to produce germanates ([GeO3]2−). Germanium occurs mostly in the oxidation state +4 although many compounds are known with the oxidation state of +2.[26] Other oxidation states are rare, such as +3 found in compounds such as Ge2Cl6, and +3 and +1 observed on the surface of oxides,[27] or negative oxidation states in germanes, such as -4 in GeH4. Germanium cluster anions (Zintl ions) such as Ge42−, Ge94−, Ge92−, [(Ge9)2]6− have been prepared by the extraction from alloys containing alkali metals and germanium in liquid ammonia in the presence of ethylenediamine or a cryptand.[26] [28] The oxidation states of the element in these ions are not integers—similar to the ozonides O3−. Two oxides of germanium are known: germanium dioxide (GeO2, germania) and germanium monoxide, (GeO).[20] The dioxide, GeO2 can be obtained by roasting germanium sulfide (GeS2), and is a white powder that is only slightly soluble in water but reacts with alkalis to form germanates.[20] The monoxide, germanous oxide, can be obtained by the high temperature reaction of GeO2 with Ge metal.[20] The dioxide (and the related oxides and germanates) exhibits the unusual property of having a high refractive index for visible light, but transparency to infrared light.[29] [30] Bismuth germanate, Bi4Ge3O12, (BGO) is used as a scintillator.[31]
5
Germanium Binary compounds with other chalcogens are also known, such as the disulfide (GeS2), diselenide (GeSe2), and the monosulfide (GeS), selenide (GeSe), and telluride (GeTe).[26] GeS2 forms as a white precipitate when hydrogen sulfide is passed through strongly acid solutions containing Ge(IV).[26] The disulfide is appreciably soluble in water and in solutions of caustic alkalis or alkaline sulfides. Nevertheless, it is not soluble in acidic water, which allowed Winkler to discover the element.[32] By heating the disulfide in a current of hydrogen, the monosulfide (GeS) is formed, which sublimes in thin plates of a dark color and metallic luster, and is soluble in solutions of the caustic alkalis.[20] Upon melting with alkaline carbonates and sulfur, germanium compounds form salts known as thiogermanates.[33] Four tetrahalides are known. Under normal conditions GeI4 is a solid, GeF4 a gas and the others volatile liquids. For example, germanium tetrachloride, GeCl4, is obtained as a colourless fuming liquid boiling at 83.1 °C by heating the metal with chlorine.[20] All the tetrahalides are readily hydrolysed to hydrated germanium dioxide.[20] GeCl4 is used in the production of organogermanium compounds.[26] All four dihalides are known Germane is similar to and in contrast to the tetrahalides are polymeric solids.[26] methane. Additionally Ge2Cl6 and some higher compounds of formula [20] GenCl2n+2 are known. The unusual compound Ge6Cl16 has been prepared that contains the Ge5Cl12 unit with a neopentane structure.[34] Germane (GeH4) is a compound similar in structure to methane. Polygermanes—compounds that are similar to alkanes—with formula GenH2n+2 containing up to five germanium atoms are known.[26] The germanes are less volatile and less reactive than their corresponding silicon analogues.[26] GeH4 reacts with alkali metals in liquid ammonia to form white crystalline MGeH3 which contain the GeH3− anion.[26] The germanium hydrohalides with one, two and three halogen atoms are colorless reactive liquids.[26] The first organogermanium compound was synthesised by Winkler in 1887; the reaction of germanium tetrachloride with diethylzinc yielded tetraethylgermane (Ge(C2H5)4).[7] Organogermanes of the type R4Ge Nucleophilic addition with an organogermanium compound (where R is an alkyl) such as tetramethylgermane (Ge(CH3)4) and tetraethylgermane are accessed through the cheapest available germanium precursor germanium tetrachloride and alkyl nucleophiles. Organic germanium hydrides such as isobutylgermane ((CH3)2CHCH2GeH3) were found to be less hazardous and may be used as a liquid substitute for toxic germane gas in semiconductor applications. Many germanium reactive intermediates are known: germyl free radicals, germylenes (similar to carbenes), and germynes (similar to carbynes).[35] [36] The organogermanium compound 2-carboxyethylgermasesquioxane was first reported in the 1970s, and for a while was used as a dietary supplement and thought to possibly have anti-tumor qualities.[37]
6
Germanium
7
Isotopes Germanium has five naturally occurring isotopes, 70Ge, 72Ge, 73Ge, 74Ge, and 76Ge. Of these, 76Ge is very slightly radioactive, decaying by double beta decay with a half-life of 1.58 × 1021 years. 74Ge is the most common isotope, having a natural abundance of approximately 36%. 76Ge is the least common with a natural abundance of approximately 7%.[38] When bombarded with alpha particles, the isotope 72Ge will generate stable 77Se, releasing high energy electrons in the process.[39] Because of this, it is used in combination with radon for nuclear batteries.[39] At least 27 radioisotopes have also been synthesized ranging in atomic mass from 58 to 89. The most stable of these is 68Ge, decaying by electron capture with a half-life of 270.95 d. The least stable is 60Ge with a half-life of 30 ms. While most of germanium's radioisotopes decay by beta decay, 61Ge and 64Ge decay by β+ delayed proton emission.[38] 84Ge through 87 Ge also have minorβ− delayed neutron emission decay paths.[38]
Natural abundance Germanium is created through stellar nucleosynthesis, mostly by the s-process in asymptotic giant branch stars. The s-process is a slow neutron capture of lighter elements inside pulsating red giant stars.[40] Germanium has been detected in the atmosphere of Jupiter[41] and in some of the most distant stars.[42] Its abundance in the Earth's crust is approximately 1.6 ppm.[43] There are only a few minerals like argyrodite, briartite, germanite, and renierite that contain appreciable amounts of germanium, but no minable deposits exist for any of them. Nonetheless, none is mined for its germanium content.[17] [44] Some zinc-copper-lead ore bodies contain enough germanium that it can be extracted from the final ore concentrate.[43] An unusual enrichment process causes a high content of germanium in some coal seams, which was discovered by Victor Mordechai Goldschmidt during a broad survey for germanium deposits.[45] [46] The highest concentration ever found was in the Hartley coal ash with up to 1.6% of germanium.[45] [46] The coal deposits near Xilinhaote, Inner Mongolia, contain an estimated 1600 tonnes of germanium.[43]
Production Worldwide production in 2006 was roughly 100 tonnes of germanium.[17] Currently, it is recovered as a by-product from sphalerite zinc ores where it is concentrated in amounts of up to 0.3%,[47] especially from sediment-hosted, massive Zn–Pb–Cu(–Ba) deposits and carbonate-hosted Zn–Pb deposits.[43] Figures for worldwide Ge reserves are not available, but in the US it is estimated to be around 500 tonnes.[43] In 2007 35% of the demand was met by recycled germanium.[43] While it is produced mainly from sphalerite, it is also found in silver, lead, and copper ores. Another source of germanium is fly ash of coal power plants which use coal from certain coal deposits with a large concentration of germanium. Russia and China Renierite
used this as a source for germanium.[48] Russia's deposits are located in the far east of the country on Sakhalin Island. The coal mines northeast of Vladivostok have also been used as
Germanium
8
a germanium source.[43] The deposits in China are mainly located in the lignite mines near Lincang, Yunnan; coal mines near Xilinhaote, Inner Mongolia are also used.[43] Year
Cost [49] ($/kg)
1999
1,400
2000
1,250
2001
890
2002
620
2003
380
2004
600
2005
660
2006
880
2007
1240
The ore concentrates are mostly sulfidic; they are converted to the oxides by heating under air, in a process known as roasting: GeS2 + 3 O2 → GeO2 + 2 SO2 Part of the germanium ends up in the dust produced during this process, while the rest is converted to germanates which are leached together with the zinc from the cinder by sulfuric acid. After neutralisation only the zinc stays in solution and the precipitate contains the germanium and other metals. After reducing the amount of zinc in the precipitate by the Waelz process, the residing Waelz oxide is leached a second time. The dioxide is obtained as precipitate and converted with chlorine gas or hydrochloric acid to germanium tetrachloride, which has a low boiling point and can be distilled off:[48] GeO2 + 4 HCl → GeCl4 + 2 H2O GeO2 + 2 Cl2 → GeCl4 + O2 Germanium tetrachloride is either hydrolysed to the oxide (GeO2) or purified by fractional distillation and then hydrolysed.[48] The highly pure GeO2 is now suitable for the production of germanium glass. The pure germanium oxide is reduced by the reaction with hydrogen to obtain germanium suitable for the infrared optics or semiconductor industry: GeO2 + 4 H2 → Ge + 2 H2O The germanium for steel production and other industrial processes is normally reduced using carbon:[50] GeO2 + C → Ge + CO2
Germanium
Applications The major end uses for germanium in 2007, worldwide, were estimated to be: 35% for fiber-optic systems, 30% infrared optics, 15% for polymerization catalysts, and 15% for electronics and solar electric [17] applications. The remaining 5% went into other uses such as phosphors, metallurgy, and chemotherapy.[17]
Optics The most notable physical characteristics of germania (GeO2) are its high index of refraction and its low optical dispersion. These make it especially useful for A typical single-mode optical fiber. Germanium oxide wide-angle camera lenses, microscopy, and is a dopant of the core silica (Item 1). 1.- Core 8 µm for the core part of optical fibers.[51] [52] It 2.Cladding 125 µm also replaced titania as the silica dopant for 3.- Buffer 250 µm silica fiber, eliminating the need for 4.- Jacket 400 µm subsequent heat treatment, which made the fibers brittle.[53] At the end of 2002 the fiber optics industry accounted for 60% of the annual germanium use in the United States, but this use accounts for less than 10% of world wide consumption.[52] GeSbTe is a phase change alloy used for its optic properties, such as in rewritable DVDs.[54] Because germanium is transparent in the infrared it is a very important infrared optical material, that can be readily cut and polished into lenses and windows. It is especially used as the front optic in thermal imaging cameras working in the 8 to 14 micron wavelength range for passive thermal imaging and for hot-spot detection in military, night vision system in cars, and fire fighting applications.[50] It is therefore used in infrared spectroscopes and other optical equipment which require extremely sensitive infrared detectors.[52]
Electronics The alloy silicon germanide (commonly referred to as "silicon-germanium", or SiGe) is rapidly becoming an important semiconductor material, for use in high speed integrated circuits. Circuits utilizing the properties of Si-SiGe junctions can be much faster than those using silicon alone.[55] Silicon-germanium is beginning to replace gallium arsenide (GaAs) in wireless communications devices.[17] The SiGe chips, with high-speed properties, can be made with low-cost, well-established production techniques of the silicon chip industry.[17] The recent rise in energy cost has improved the economics of solar panels, a potential major new use of germanium.[17] Germanium is the substrate of the wafers for high-efficiency multijunction photovoltaic cells for space applications.
9
Germanium
10
Gallium arsenide germanium solar cell Because germanium and gallium arsenide have very similar lattice constants, germanium substrates can be used to make gallium arsenide solar cells.[56] The Mars Exploration Rovers and several satellites use triple junction gallium arsenide on germanium cells.[57] Germanium-on-insulator substrates are seen as a potential replacement for silicon on miniaturized chips.[17] Other uses in electronics include phosphors in fluorescent lamps,[21] and germanium-base solid-state light-emitting diodes (LEDs).[17] Germanium transistors are still used in some effects pedals by musicians who wish to reproduce the distinctive tonal character of the "fuzz"-tone from the early rock and roll era, most notably the Dallas Arbiter Fuzz Face.[58]
Other uses Germanium
dioxide
is
also
used
in
catalysts
for
polymerisation in the production of polyethylene [59] terephthalate (PET). The high brilliance of the produced polyester is especially used for PET bottles marketed in Japan.[59] However, in the United States, no germanium is used for polymerization catalysts.[17] Due to the similarity between silica (SiO2) and germanium dioxide (GeO2), the silica stationary phase in some gas chromatography columns can be replaced by GeO2.[60] In recent years germanium has seen increasing use in precious metal alloys. In sterling silver alloys, for instance, it has been found to reduce firescale, increase tarnish resistance, and increase the alloy's response to precipitation hardening. A tarnish-proof sterling silver alloy, trademarked A PET bottle Argentium, requires 1.2% germanium.[17] The material has a very high refractive index (4.0) and so needs to be anti-reflection coated. Particularly, a very hard special antireflection coating of diamond-like carbon (DLC), refractive index 2.0, is a good match and produces a diamond-hard surface that can withstand much environmental rough treatment.[61] [62] High purity germanium single crystal detectors can precisely identify radiation sources—for example in airport security.[63] Germanium is useful for monochromators for beamlines used in single crystal neutron scattering and synchrotron X-ray diffraction. The reflectivity has advantages over silicon in neutron and high energy X-ray applications.[64] Crystals of high purity germanium are used in detectors for gamma spectroscopy and the search for dark matter.[65] Certain compounds of germanium have low toxicity to mammals, but have toxic effects against certain bacteria.[19] This property makes these compounds useful as chemotherapeutic agents.[66]
Germanium
11
Precautions As early as 1922, doctors in the United States used the inorganic form of germanium to treat patients with anemia.[67] It was used in other forms of treatments, but its efficiency has been dubious. Its role in cancer treatments has been debated.[68] U.S. Food and Drug Administration research has concluded that germanium, when used as a nutritional supplement, "presents potential human health hazard".[37] Germanium is not thought to be essential to the health of plants or animals. Some of its compounds present a hazard to human health, however. For example, germanium chloride and germane (GeH4) are a liquid and gas, respectively, that can be very irritating to the eyes, skin, lungs, and throat.[69] Germanium has little or no impact on the environment because it usually occurs only as a trace element in ores and carbonaceous materials, and is used in very small quantities in commercial applications.[17]
Footnotes 1.
^
From Greek, argyrodite means silver-containing.[70]
2.
^
Just as the existence of the new element was predicted, the existence of the planet Neptune was predicted around 1843 by the mathematicians John Couch Adams and Urbain Leverrier for the fact that Uranus was being pulled slightly out of position in its orbit.[71] James Challis started searching for it in July 1846 and sighted the planet 23 September 1846.[72] 3. ^ R. Hermann published in 1877 claims of the discovery of a new element beneath tantalum, which he named neptunium.[73] [74] But this was later regarded as some mixture of niobium and tantalum.[75] The name neptunium was eventually given to the synthetic element past uranium discovered in 1940.[76]
External links • WebElements.com – Germanium
[77]
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"Germanium phoenix". III-Vs Review 18 (7): 55. doi: 10.1016/S0961-1290(05)71310-7 (http:/ / dx. doi. org/ 10. 1016/ S0961-1290(05)71310-7). [59] Thiele, Ulrich K. (2001). "The Current Status of Catalysis and Catalyst Development for the Industrial Process of Poly(ethylene terephthalate) Polycondensation". International Journal of Polymeric Materials 50 (3): 387 – 394. doi: 10.1080/00914030108035115 (http:/ / dx. doi. org/ 10. 1080/ 00914030108035115). [60] Fang, Li; Kulkarni, Sameer; Alhooshani, Khalid; Malik, Abdul (2007). "Germania-Based, Sol-Gel Hybrid Organic-Inorganic Coatings for Capillary Microextraction and Gas Chromatography". Anal. Chem. 79 (24): 9441–9451. doi: 10.1021/ac071056f (http:/ / dx. doi. org/ 10. 1021/ ac071056f). [61] Lettington, Alan H. (1998). "Applications of diamond-like carbon thin films". Carbon 36 (5–6): 555–560. doi: 10.1016/S0008-6223(98)00062-1 (http:/ / dx. doi. org/ 10. 1016/ S0008-6223(98)00062-1). [62] Gardos, Michael N.; Bonnie L. Soriano, Steven H. Propst (1990). "Study on correlating rain erosion resistance with sliding abrasion resistance of DLC on germanium". Proc. SPIE, 1325 (Mechanical Properties): 99. doi: 10.1117/12.22449 (http:/ / dx. doi. org/ 10. 1117/ 12. 22449). [63] Keyser, Ronald. " Performance of Light-Weight, Battery-Operated, High Purity Germanium Detectors for Field Use (http:/ / www. ortec-online. com/ papers/ inmm_2003_keyser. pdf)" (pdf). Oak Ridge Technical Enterprise Corporation (ORTEC). . Retrieved 2008-09-06. [64] Ahmed, F. U. (1996). "Optimization of Germanium for Neutron Diffractometers". International Journal of Modern Physics E 5: 131. doi: 10.1142/S0218301396000062 (http:/ / dx. doi. org/ 10. 1142/ S0218301396000062). [65] Diehl, R. (2006). "Astrophysical constraints from gamma-ray spectroscopy". Nuclear Physics A 777: 70. doi: 10.1016/j.nuclphysa.2005.02.155 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2005. 02. 155). [66] Slavik, Milan; Blanc, Oscar; Davis, Joan (1983). "Spirogermanium: A new investigational drug of novel structure and lack of bone marrow toxicity". Investigational New Drugs 1 (3): 225–234. doi: 10.1007/BF00208894 (http:/ / dx. doi. org/ 10. 1007/ BF00208894). [67] Brown Jr., Robert D. Commodity Survey:Germanium (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ germanium/ 220798. pdf), US Geological Surveys. Report. Retrieved on 2008-09-09. [68] " Germanium (http:/ / www. cancer. org/ docroot/ ETO/ content/ ETO_5_3X_Germanium. asp)". American Cancer Society. . Retrieved 2008-08-31. [69] Gerber, G.B.; Léonard, A. (1997). "Mutagenicity, carcinogenicity and teratogenicity of germanium compounds". Regulatory Toxicology and Pharmacology 387: 141–146. doi: 10.1016/S1383-5742(97)00034-3 (http:/ / dx. doi. org/ 10. 1016/ S1383-5742(97)00034-3). [70] Argyrodite—Ag8GeS6 (http:/ / www. handbookofmineralogy. org/ pdfs/ argyrodite. pdf), Mineral Data Publishing. Report. Retrieved on 2008-09-01. [71] Adams, J. C. (November 13, 1846). " Explanation of the observed irregularities in the motion of Uranus, on the hypothesis of disturbance by a more distant planet (http:/ / adsabs. harvard. edu/ cgi-bin/ nph-bib_query?bibcode=1846MNRAS. . . 7. . 149A& amp;db_key=AST& amp;data_type=HTML& amp;format=& amp;high=42c888df4622238)". Monthly Notices of the Royal Astronomical Society (Blackwell Publishing) 7: 149. . Retrieved 2008-02-18. [72] Challis, Rev. J. (November 13, 1846). " Account of observations at the Cambridge observatory for detecting the planet exterior to Uranus (http:/ / adsabs. harvard. edu/ cgi-bin/ nph-bib_query?bibcode=1846MNRAS. . . 7. . 145C& amp;db_key=AST& amp;data_type=HTML& amp;format=& amp;high=42c888df4622238)". Monthly Notices of the Royal Astronomical Society (Blackwell Publishing) 7: 145–149. . Retrieved 2008-02-18. [73] Sears, Robert (July 1877). "Scientific Miscellany". The Galaxy (Columbus, O[hio]: Siebert & Lilley) 24 (1): 131. ISBN 0665501668. OCLC 16890343 243523661 77121148 (http:/ / worldcat. org/ oclc/ 16890343+ 243523661+ 77121148). [74] " Editor's Scientific Record (http:/ / cdl. library. cornell. edu/ cgi-bin/ moa/ moa-cgi?notisid=ABK4014-0055-21)". Harper's new monthly magazine 55 (325): 152–153. June 1877. . [75] van der Krogt, Peter. " Elementymology & Elements Multidict: Niobium (http:/ / elements. vanderkrogt. net/ elem/ nb. html)". . Retrieved 2008-08-20.
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Germanium [76] Westgren, A. (1964). " The Nobel Prize in Chemistry 1951: presentation speech (http:/ / nobelprize. org/ nobel_prizes/ chemistry/ laureates/ 1951/ press. html)". Nobel Lectures, Chemistry 1942-1962. Elsevier. . [77] http:/ / www. webelements. com/ webelements/ elements/ text/ Ge/ index. html
15
Article Sources and Contributors
Article Sources and Contributors Germanium Source: http://en.wikipedia.org/w/index.php?oldid=306030695 Contributors: ABF, AJim, Abe Lincoln, Acolorpink1, Acroterion, Ahoerstemeier, Aitias, Aka042, Alexdeh, Amalthea, Anclation, Archimerged, Art LaPella, AssegaiAli, Avnjay, Axiosaurus, B.d.mills, Badinfinity, Barathrumm, Barticus88, Bdesham, Beano, Beetstra, Benbest, Berkay0652, Biscuittin, Blanchardb, BlueEarth, BlytheG, Bongwarrior, Bovineone, Bryan Derksen, Buchanan-Hermit, Butwhatdoiknow, CYD, Can't sleep, clown will eat me, CanisRufus, Capricorn42, Captainbeefart, Carnildo, ChemGardener, Chris 73, Commander Shepard, Conversion script, Cryptic C62, DMacks, Darrien, David Latapie, Deli nk, Delldot, Delta G, Digitalfear, Discospinster, Dlohcierekim's sock, Donarreiskoffer, Doodle77, DrBob, Dshenai, E Wing, EchetusXe, Edgar181, Edwy, El C, Eleuther, Eloil, Ember of Light, Emperorbma, Epbr123, Erik Zachte, FelisSchrödingeris, Femto, FlyingToaster, Frankenpuppy, Fromjarod, Galoubet, Gdarin, Geoffrey.landis, Giftlite, Giggy, Gman124, Graham87, GrahamHardy, Graue, Grendelkhan, Ground Zero, Gscshoyru, Hadal, Hak-kâ-ngìn, Hcaz11, Headbomb, Hedgerhedger, Hellbus, Hgjghjh, Hydrargyrum, II MusLiM HyBRiD II, Icairns, Igoldste, Intelligentguest, Irishleprechaun, Itub, J.delanoy, Jaan513, Jackol, Janke, Jaraalbe, Jerzy, Jimfbleak, Joanjoc, John, Jopusbob, Jose77, Josh Grosse, Joyous!, Juanathan, KJS77, Kablammo, Kaimiddleton, Karlhahn, Kdliss, Keilana, Keithpoole, King of Hearts, Knepflerle, Krellis, Kurykh, Kwamikagami, LA2, LarryMorseDCOhio, Legoktm, Lethalgeek, LeyteWolfer, Linas, Ling.Nut, Loveandprotect, Lugh23, LuigiManiac, Luk, Mac, Magog the Ogre, MarsRover, Materialscientist, Mav, Mgimpel, Minesweeper, Mr0t1633, Mukkakukaku, NeonMerlin, Nergaal, New4325, NewEnglandYankee, Nibuod, Nihiltres, Novangelis, Novel tubes, Nvj, OldZeb, Opelio, Ottava Rima, Ottre, Oudegeest, Paul August, Persian Poet Gal, PeterJeremy, PhazZ, Philip Trueman, PierreAbbat, Piledhigheranddeeper, Plasmic Physics, Plexust, Poolkris, Poor Yorick, Potatoswatter, PranksterTurtle, PsychoCola, Qleem, Qxz, R. fiend, R9tgokunks, RJHall, RTC, Remember, Retiono Virginian, RexxS, Rjwilmsi, Roberta F., Romanm, Rrburke, Runch, SEWilco, SaltyBoatr, Saperaud, Schneelocke, Scorpion451, Scott940603, Sengkang, Sergio.ballestrero, Shaddack, Shellreef, Sifaka, Sl, Smokefoot, Somearemoreequal, Squids and Chips, Srleffler, StaticGull, Stifynsemons, Stone, StradivariusTV, Suisui, Sunborn, Tantalate, Tempodivalse, Tetracube, The Hemp Necktie, The Ogre, The Red, Thincat, Thingg, Thricecube, Thue, Tide rolls, Tim Starling, TravisAF, Trevas, Trevor MacInnis, Trevor f, Tylerdmace, Verkhovensky, Versus22, Vsmith, WCFrancis, Warut, Watch37264, Wavelength, WereSpielChequers, Wik1ped1a is meant 2 be vanda1ised, WillMak050389, Willking1979, WolfmanSF, XLerate, Yekrats, Yyy, Zubaexy, 373 anonymous edits
Image Sources, Licenses and Contributors image:Ge-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Ge-TableImage.png License: GNU Free Documentation License Contributors: user:Schneelocke Image: Germanium.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Germanium.jpg License: GNU Free Documentation License Contributors: Kanonkas, Saperaud Image:medeleeff by repin.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Medeleeff_by_repin.jpg License: unknown Contributors: Hailey C. Shannon, J.M.Domingo, Kevyn, Kneiphof, Maximaximax, OldakQuill, Ragesoss, Solon, XJamRastafire, 竹麦魚(Searobin), 2 anonymous edits Image:Winkler Clemens.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Winkler_Clemens.jpg License: Public Domain Contributors: Astrochemist Image:Germane-2D-dimensions.png Source: http://en.wikipedia.org/w/index.php?title=File:Germane-2D-dimensions.png License: Public Domain Contributors: Ben Mills Image:NucleophilicAdditionWithOrganogermanium.png Source: http://en.wikipedia.org/w/index.php?title=File:NucleophilicAdditionWithOrganogermanium.png License: GNU Free Documentation License Contributors: Mahahahaneapneap, V8rik Image:Renierit.JPG Source: http://en.wikipedia.org/w/index.php?title=File:Renierit.JPG License: Creative Commons Attribution-Sharealike 2.0 Contributors: Alchemist-hp (www.pse-mendelejew.de) Original uploader was Alchemist-hp at de.wikipedia File:Singlemode fibre structure.svg Source: http://en.wikipedia.org/w/index.php?title=File:Singlemode_fibre_structure.svg License: unknown Contributors: User:Benchill, User:DrBob Image:Pet Flasche.JPG Source: http://en.wikipedia.org/w/index.php?title=File:Pet_Flasche.JPG License: unknown Contributors: Ttox, 1 anonymous edits
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
16
Arsenic
1
Arsenic germanium ← arsenic → seleniumP ↑ As ↓ Sb
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
33As Periodic table
Appearance metallic grey
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Sublimation pointTriple pointCritical pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
553
596
646
706
781
874
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1798 kJ·mol−1 3rd: 2735 kJ·mol−1Atomic radiusCovalent radiusVan der Waals
Arsenic
2
radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityYoung's modulusBulk modulusMohs hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of arsenic iso 73
As
74
As
75
As
N.A. syn
syn
100%
half-life 80.3 d
17.78 d
DM
DE (MeV)
DP
ε
-
73
γ
0.05D, 0.01D, e
-
ε
-
74
β+
0.941
74
γ
0.595, 0.634
-
β−
1.35, 0.717
74
Ge
Ge Ge
Se
75
As is stable with 42 neutron
arsenic, As, 33 metalloid15, 4, p74.92160(2) g·mol−1 [Ar] 4s2 3d10 4p3 2, 8, 18, 5 (Image) solid 5.727 g·cm−3 5.22 g·cm−3 887 K,615 °C,1139 °F 1090 K (817°C), 3628 [1] kPa 1673 K, ? MPa (grey) 24.44 kJ·mol−1 ? 34.76 kJ·mol−1 (25 °C) 24.64 J·mol−1·K−15, 3, 2, 1,[2] -3 (mildly acidic oxide) 2.18 (Pauling scale) 1st: 947.0 kJ·mol−1119 pm119±4 pm 185 pm rhombohedral diamagnetic[3] (20 °C) 333 nΩ·m (300 K) 50.2 W·m−1·K−1 8 GPa 22 GPa 3.5 1440 MPa 7440-38-2 Arsenic (pronounced /ˈɑrsnɪk/; also English pronunciation: /ɑrˈsɛnɪk/ when attributive) is the chemical element that has the symbol As and atomic number 33. Arsenic was first documented by Albertus Magnus in 1250.[4] Its atomic mass is 74.92. Arsenic is a notoriously poisonous metalloid with many allotropic forms, including a yellow (molecular non-metallic) and several black and grey forms (metalloids). Three metalloidal forms of arsenic, each with a different crystal structure, are found free in nature (the minerals arsenic sensu stricto and the much rarer arsenolamprite and pararsenolamprite). However, it is more commonly found as arsenide and in arsenate compounds, several hundred of which are known. Arsenic and its compounds are used as pesticides, herbicides, insecticides and in various alloys.
History The word arsenic was borrowed from the Persian word خينرزZarnikh, meaning "yellow orpiment", into Greek as arsenikon. It is also related to the similar Greek word "arsenikos", meaning "masculine" or "potent". Arsenic compounds (orpiment, realgar) have been known and used since ancient times.[5] As the symptoms of arsenic poisoning were somewhat ill-defined, it was frequently used for murder until the advent of the Marsh test, a sensitive chemical test for its presence. (Another less sensitive but more general test is the Reinsch test.) Due to its use by the ruling class to murder one another and its potency and discreetness, arsenic has been called the Poison of Kings and the King of Poisons.[6] During the Bronze Age, arsenic was often included in bronze, which made the alloy harder (so-called "arsenical bronze"). Albertus Magnus (Albert the Great, 1193–1280) is believed to have been the first European to isolate the element in 1250 by heating soap together with arsenic trisulfide.[4] In 1649, Johann Schröder published two ways of preparing arsenic.
Arsenic
3 Cadet's fuming liquid (impure cacodyl), the first organometallic compound, was synthesized in 1760 by Louis Claude Cadet de Gassicourt by the reaction of potassium acetate with arsenic trioxide.[7]
In the Victorian era, "arsenic" (colourless, crystalline, soluble "white arsenic" trioxide) was mixed with vinegar and chalk and eaten by women Alchemical symbol for to improve the complexion of their faces, making their skin paler to show arsenic they did not work in the fields. Arsenic was also rubbed into the faces and arms of women to "improve their complexion". The accidental use of arsenic in the adulteration of foodstuffs led to the Bradford sweet poisoning in 1858, which resulted in approximately 20 deaths and 200 people taken ill with arsenic poisoning.[8]
Characteristics Isotopes Naturally occurring arsenic is composed of one stable isotope, 75As.[9] As of 2003, at least 33 radioisotopes have also been synthesized, ranging in atomic mass from 60 to 92. The most stable of these is 73As with a half-life of 80.3 days. Isotopes that are lighter than the stable 75As tend to decay by β+ decay, and those that are heavier tend to decay by β- decay, with some exceptions. At least 10 nuclear isomers have been described, ranging in atomic mass from 66 to 84. The most stable of arsenic's isomers is 68mAs with a half-life of 111 seconds.[9]
Allotropes Like phosphorus, arsenic is an excellent example of an element that exhibits allotropy, as its various allotropes have strikingly different properties. The three most common allotropes are metallic grey, yellow and black arsenic.[10] The most common allotrope of arsenic is grey arsenic. It has a similar structure to black phosphorus (β-metallic phosphorus) and has a layered crystal structure somewhat resembling that of graphite. It consists of many six-membered rings which are interlinked. Each atom is bound to three other atoms in the layer and is coordinated by each 3 arsenic atoms in the upper and lower layer. This relatively close packing leads to a high density of 5.73 g/cm3.[11] Structure of yellow arsenic As4 and white phosphorus P4
Yellow arsenic (As4) is soft and waxy, not unlike P4. Both have four atoms arranged in a tetrahedral structure in which each atom is bound to the other three atoms by a single bond, resulting in very high ring strain and instability. This form of the elements are the least stable, most reactive, more volatile, less dense, and more toxic than the other allotropes. Yellow arsenic is produced by rapid cooling of arsenic vapour with liquid nitrogen. It is rapidly transformed into the grey arsenic by light. The yellow form has a density of 1.97 g/cm³.[11] Black arsenic is similar in structure to red phosphorus.[11]
Arsenic
Chemical The most common oxidation states for arsenic are −3 (arsenides: usually alloy-like intermetallic compounds), +3 (arsenates(III) or arsenites, and most organoarsenic compounds), and +5 (arsenates: the most stable inorganic arsenic oxycompounds). Arsenic also bonds readily to itself, forming square As3−4 ions in the arsenide skutterudite. In the +3 oxidation state, the stereochemistry of arsenic is affected by possession of a lone pair of electrons. Arsenic is very similar chemically to its predecessor in the Periodic Table, phosphorus. Like phosphorus, it forms colourless, odourless, crystalline oxides As2O3 and As2O5 which are hygroscopic and readily soluble in water to form acidic solutions. Arsenic(V) acid is a weak acid. Like phosphorus, arsenic forms an unstable, gaseous hydride: arsine (AsH3). The similarity is so great that arsenic will partly substitute for phosphorus in biochemical reactions and is thus poisonous. However, in subtoxic doses, soluble arsenic compounds act as stimulants, and were once popular in small doses as medicinals by people in the mid 18th century.[12] When heated in air, arsenic oxidizes to arsenic trioxide; the fumes from this reaction have an odour resembling garlic. This odour can be detected on striking arsenide minerals such as arsenopyrite with a hammer. Arsenic (and some arsenic compounds) sublimes upon heating at atmospheric pressure, converting directly to a gaseous form without an intervening liquid state. The liquid state appears at 20 atmospheres and above, which explains why the melting point is higher than the boiling point.[11]
Compounds Arsenic compounds resemble in many respects those of phosphorus as both arsenic and phosphorus occur in the same group (column) of the periodic table. The most important compounds of arsenic are arsenic(III) oxide, As2O3, ("white arsenic"), the yellow sulfide orpiment (As2S3) and red realgar (As4S4), Paris Green, calcium arsenate, and lead hydrogen arsenate. The latter three have been used as agricultural insecticides and poisons. Whilst arsenic trioxide forms during oxidation of arsenic, arsenic pentoxide is formed by the dehydration of arsenic acid. Both oxides dissolve in strong alkaline solution, with the formation of arsenite AsO3−3 and arsenate AsO3−4 respectively. The protonation steps between the arsenate and arsenic acid are similar to those between phosphate and phosphoric acid. However, arsenite and arsenous acid contain arsenic bonded to three oxygens and not hydrogen, in contrast to phosphite and phosphorous acid (more accurately termed 'phosphonic acid'), which contain non-acidic P-H bonds. Arsenous acid is genuinely tribasic, whereas phosphonic acid is not. A broad variety of sulfur compounds of arsenic are known, As4S3, As4S4, As2S3 and As4S10. All arsenic(III) halogen compounds (except with astatine) are known and stable. For the arsenic(V) compounds the situation is different: only the arsenic pentafluoride is stable at room temperature. Arsenic pentachloride is only stable at temperatures below -50 °C and the pentabromide and pentaiodide are unknown.[12] Arsenic is used as group 5 element as part of the III-V semiconducting compounds. Gallium arsenide, indium arsenide and aluminium arsenide are used as semiconductor material when the properties of silicon are not suitable for the application and the higher price of
4
Arsenic
5
the compounds is acceptable. • • • • • • •
Arsenic acid (H3AsO4) Arsenous acid (H3AsO3) Arsenic trioxide (As2O3) Arsine (Arsenic Trihydride AsH3) Cadmium arsenide (Cd3As2) Gallium arsenide (GaAs) Lead hydrogen arsenate (PbHAsO4)
Arsenic also has a formal oxidation state of +2 in As4S4, realgar. This is achieved by pairing As atoms to produce dimeric cations [As-As]2+, so the total covalency of As is still in fact three. [13]
Occurrence Arsenopyrite, also unofficially called mispickel,[14] (FeAsS) is the most common arsenic-bearing mineral. In the litosphere the minerals of the formula M(II)AsS, with M(II) being mostly Fe, Ni and Co, are the dominat arsenic minerals.
A large sample of native arsenic.
Orpiment and realgar were formerly used as painting pigments, though they have fallen out of use due to their toxicity and reactivity. Although arsenic is sometimes found native in nature, its main economic source is the mineral arsenopyrite mentioned above; it is also found in arsenides of metals such as silver, cobalt (cobaltite: CoAsS and skutterudite: CoAs3) and Realgar nickel, as sulfides, and when oxidised as arsenate minerals such as mimetite, Pb5(AsO4)3Cl and erythrite, Co3(AsO4)2. 8H2O, and more rarely arsenites ('arsenite' = arsenate(III), AsO33- as opposed to arsenate (V), AsO43-). In addition to the inorganic forms mentioned above, arsenic also occurs in various organic forms in the environment. [15]
Arsenic
Production In 2005, China was the top producer of white arsenic with almost 50% world share, followed by Chile, Peru and Morocco, reports the British Geological Survey and the United States Geological Survey.[16] The arsenic was recovered mostly during mining operations, for example the [16] production from Peru comes mostly from Arsenic output in 2006 copper mining and the production in China is due to gold mining. Arsenic is part of the smelter dust from copper, gold, and lead smelters.[17] On roasting in air of arsenopyrite, arsenic sublimes as arsenic (III) oxide leaving iron oxides[15] , while roasting without air results in the production of metallic arsenic. For further purification of the arsenic from sulfur and other chacogenes is sublimed in vacuum or in a hydrogen atmosphere or by distilling it from molten lead-arsenic mixture.[18]
Applications Wood preservation The toxicity of arsenic to insects, bacteria and fungi makes it an ideal component for the preservation of wood. The worldwide treatment with chromated copper arsenate, also known as CCA or Tanalith was the largest consumer of arsenic since the introduction of the process in the 1950s. Due to the environmental problems caused by the arsenic most countries banned the use of chromated copper arsenate on consumer products. The ban began in the European Union and in the United States in 2004.[19] [20] In 2002 in the United States 90% of the 19,600 metric tons of arsenic compounds were used to preserve wood, in 2007 still 50% of the 5,280 metric tons of consumption was used for this purpose.[17] [21] In the European Union the use of arsenic in consumer products According to the USEPA's website, CCA lumber was discontinued for residential and general consumer construction on December 31, 2003 and alternative methods are now used like ACQ, Borates, Copper Azole, Cyproconazole, and Propiconazole. Although discontinued, this application is also one of the most concern to the general public. The vast majority of older pressure-treated wood was treated with CCA. CCA lumber is still in widespread use in many countries, and was heavily used during the latter half of the 20th century as a structural and outdoor building material. Although the use of CCA lumber was banned in many areas after studies showed that arsenic could leach out of the wood into the surrounding soil (from playground equipment, for instance), a risk is also presented by the burning of older CCA timber. The direct or indirect ingestion of wood ash from burnt CCA lumber has caused fatalities in animals and serious poisonings in humans; the lethal human dose is approximately 20 grams of ash. Scrap CCA lumber from construction and demolition sites may be inadvertently used in commercial and domestic fires. Protocols for safe disposal of CCA lumber do not exist evenly throughout the world; there is also concern in some quarters about the widespread landfill disposal of such timber.[22]
6
Arsenic
Medical During the 18th, 19th, and 20th centuries, a number of arsenic compounds have been used as medicines, including arsphenamine (by Paul Ehrlich) and arsenic trioxide (by Thomas Fowler). Arsphenamine as well as Neosalvarsan was indicated for syphilis and trypanosomiasis, but has been superseded by modern Arsphenamine (Salvarsan) antibiotics. Arsenic trioxide has been used in a variety of ways over the past 500 years, but most commonly in the treatment of cancer. The US Food and Drug Administration in 2000 approved this compound for the treatment of patients with acute promyelocytic leukemia that is resistant to ATRA.[23] It was also used as Fowler's solution in psoriasis.[24] Recently new research has been done in locating tumours using arsenic-74 (a positron emitter). The advantages of using this isotope instead of the previously used iodine-124 is that the signal in the PET scan is clearer as the iodine tends to transport iodine to the thyroid gland producing a lot of noise.[25]
Pigments Copper acetoarsenite was used as a green pigment known under many different names, including 'Paris Green' and 'Emerald Green'. It caused numerous arsenic poisonings. Scheele's Green, a copper arsenate, was used in the 19th century as a coloring agent in sweets.[26]
Military After World War I the United States built up a stockpile of 20,000 tons of Lewisite; a chemical weapon, acting as a vesicant (blister agent) and lung irritant. The stockpile was neutralized with bleach and dumped into the Gulf of Mexico after the 1950s.[27] During the Vietnam War the United States used Agent Blue (a mixture of sodium cacodylate) and dimethyl arsinic acid (cacodylic acid) as one of the rainbow herbicides to deprive the Vietnamese of valuable crops.
Other uses • Various agricultural insecticides, termination and poisons. For example Lead hydrogen arsenate was used well into the 20th century as an insecticide on fruit trees.[28] Its use sometimes resulted in brain damage to those working the sprayers. In the last half century, monosodium methyl arsenate (MSMA) and disodium methyl arsenate (DSMA), a less toxic organic form of arsenic, has replaced lead arsenate's role in agriculture. • Used in animal feed, particularly in the US as a method of disease prevention[29] (section 5.3, pg. 310) [30] and growth stimulation. One example is roxarsone which was used by 69.8 and 73.9% of the broiler starter and growers between 1995 to 2000.[31] • Gallium arsenide is an important semiconductor material, used in integrated circuits. Circuits made using the compound are much faster (but also much more expensive) than those made in silicon. Unlike silicon it is direct bandgap, and so can be used in laser diodes and LEDs to directly convert electricity into light. • Also used in bronzing and pyrotechnics. • Up to 2% of arsenic is used in lead alloys for lead shots and bullets.[32]
7
Arsenic
8
• Arsenic is added in small quantities to brass to make it Dezincification resistant. This grade of brass is used to make plumbing fittings. • Arsenic is also used for taxonomic sample preservation.
Biological role
Trimethylarsine
Inorganic arsenic and its compounds, upon entering the food chain, are progressively metabolised to less toxic forms of arsenic through a process of methylation. For example, the mold Scopulariopsis brevicaulis produce significant amounts of trimethylarsine if inorganic arsenic is present.[33] The organic compound arsenobetaine is found in some marine foods such as fish and algae, and also in mushrooms in larger concentrations. The average person's intake is about 10–50 µg/day. Values about 1000 µg are not unusual following consumption of fish or mushrooms. But there is little danger in eating fish because this arsenic compound is nearly non-toxic.[34] Some species of bacteria obtain their energy by
Arsenobetaine
oxidizing various fuels while reducing arsenate to arsenite. The enzymes involved are known as arsenate reductases (Arr).
In 2008, bacteria were discovered that employ a version of photosynthesis in the absence of oxygen with arsenites as electron donors, producing arsenates (just like ordinary photosynthesis uses water as electron donor, producing molecular oxygen). Researchers conjecture that historically these photosynthesizing organisms produced the arsenates that allowed the arsenate-reducing bacteria to thrive. One strain PHS-1 has been isolated and is related to the γ-Proteobacterium Ectothiorhodospira shaposhnikovii. The mechanism is unknown, but an encoded Arr enzyme may function in reverse to its known homologues.[35]
Safety Arsenic and many of its compounds are especially potent poisons. Arsenic disrupts ATP production through several mechanisms. At the level of the citric acid cycle, arsenic inhibits pyruvate dehydrogenase and by competing with phosphate it uncouples oxidative phosphorylation, thus inhibiting energy-linked reduction of NAD+, mitochondrial respiration, and ATP synthesis. Hydrogen peroxide production is also increased, which might form reactive oxygen species and oxidative stress. These metabolic interferences lead to death from multi-system organ failure, probably from necrotic cell death, not apoptosis. A post mortem reveals brick red colored mucosa, due to severe hemorrhage. Although arsenic causes toxicity, it can also play a protective role.[36] Elemental arsenic and arsenic compounds are classified as "toxic" and "dangerous for the environment" in the European Union under directive 67/548/EEC. The International Agency
Arsenic for Research on Cancer (IARC) recognizes arsenic and arsenic compounds as group 1 carcinogens, and the EU lists arsenic trioxide, arsenic pentoxide and arsenate salts as category 1 carcinogens. Arsenic is known to cause arsenicosis due to its manifestation in drinking water, “the most common species being arsenate [HAsO42- ; As(V)] and arsenite [H3AsO3 ; As(III)]”. The ability of arsenic to undergo redox conversion between As(III) and As(V) makes its availability in the environment more abundant. According to Croal, Gralnick, Malasarn, and Newman, “[the] understanding [of] what stimulates As(III) oxidation and/or limits As(V) reduction is relevant for bioremediation of contaminated sites (Croal). The study of chemolithoautotrophic As(III) oxidizers and the heterotrophic As(V) reducers can help the understanding of the oxidation and/or reduction of arsenic.[37] Treatment of chronic arsenic poisoning is easily accomplished. British anti-Lewisite (dimercaprol) is prescribed in dosages of 5 mg/kg up to 300 mg each 4 hours for the first day. Then administer the same dosage each 6 hours for the second day. Then prescribe this dosage each 8 hours for eight additional days.[38]
Arsenic in drinking water Arsenic contamination of groundwater has led to a massive epidemic of arsenic poisoning in Bangladesh[39] and neighbouring countries. Presently 42 major incidents around the world have been reported on groundwater arsenic contamination. It is estimated that approximately 57 million people are drinking groundwater with arsenic concentrations elevated above the World Health Organization's standard of 10 parts per billion. However, a study of cancer rates in Taiwan[40] suggested that significant increases in cancer mortality appear only at levels above 150 parts per billion. The arsenic in the groundwater is of natural origin, and is released from the sediment into the groundwater due to the anoxic conditions of the subsurface. This groundwater began to be used after local and western NGOs and the Bangladeshi government undertook a massive shallow tube well drinking-water program in the late twentieth century. This program was designed to prevent drinking of bacterially contaminated surface waters, but failed to test for arsenic in the groundwater. Many other countries and districts in South East Asia, such as Vietnam, Cambodia, and China have geological environments conducive to generation of high-arsenic groundwaters. Arsenicosis was reported in Nakhon Si Thammarat, Thailand in 1987, and the dissolved arsenic in the Chao Phraya River is suspected of containing high levels of naturally occurring arsenic, but has not been a public health problem due to the use of bottled water.[41] The northern United States, including parts of Michigan, Wisconsin, Minnesota and the Dakotas are known to have significant concentrations of arsenic in ground water. Increased levels of skin cancer have been associated with arsenic exposure in Wisconsin, even at levels below the 10 part per billion drinking water standard.[42] Epidemiological evidence from Chile shows a dose dependent connection between chronic arsenic exposure and various forms of cancer, particularly when other risk factors, such as cigarette smoking, are present. These effects have been demonstrated to persist below 50 parts per billion.[43] Analyzing multiple epidemiological studies on inorganic arsenic exposure suggests a small but measurable risk increase for bladder cancer at 10 parts per billion.[44] According to Peter Ravenscroft of the Department of Geography at the University of Cambridge,[45]
9
Arsenic roughly 80 million people worldwide consume between 10 and 50 parts per billion arsenic in their drinking water. If they all consumed exactly 10 parts per billion arsenic in their drinking water, the previously cited multiple epidemiological study analysis would predict an additional 2,000 cases of bladder cancer alone. This represents a clear underestimate of the overall impact, since it does not include lung or skin cancer, and explicitly underestimates the exposure. Those exposed to levels of arsenic above the current WHO standard should weigh the costs and benefits of arsenic remediation. Arsenic can be removed from drinking water through coprecipitation of iron minerals by oxidation and filtering. When this treatment fails to produce acceptable results, adsorptive arsenic removal media may be utilized. Several adsorptive media systems have been approved for point-of-service use in a study funded by the United States Environmental Protection Agency (U.S.EPA) and the National Science Foundation (NSF). A team of European and Indian Engineers led by Dr. Bhaskar Sen Gupta of Queen's University Belfast have set up six arsenic treatment plants in West Bengal based on in-situ remediation method. This technology does not use any chemicals and arsenic is turned into an insoluble form in the subterranean zone by recharging aerated water into the aquifer. This process does not produce any waste stream or sludge and it costs about 1 US$ to produce 10 cubic meters of water [46] Magnetic separations of arsenic at very low magnetic field gradients have been demonstrated in point-of-use water purification with high-surface-area and monodisperse magnetite (Fe3O4) nanocrystals. Using the high specific surface area of Fe3O4 nanocrystals the mass of waste associated with arsenic removal from water has been dramatically reduced.[47] Epidemiological studies have suggested a correlation between chronic consumption of drinking water contaminated with arsenic and the incidence of type 2 diabetes. However, the literature provides insufficient scientific evidence to show cause and effect between arsenic and the onset of diabetes mellitus type 2.
Occupational Exposures Industries that use inorganic arsenic and its compounds include wood preservation, glass production, nonferrous metal alloys, and electronic semiconductor manufacturing. Inorganic arsenic is also found in coke oven emissions associated with the smelter industry.[48] Occupational exposure and poisoning may occur in persons working in these industries.
See also • • • • • •
Aqua Tofana Arsenic poisoning Fowler's solution Grainger challenge White arsenic Arsenic trioxide
10
Arsenic
11
External links • Los Alamos National Laboratory – Arsenic
[49]
• CTD's Arsenic page [50] and CTD's Arsenicals page [51] from the Comparative Toxicogenomics Database • A Small Dose of Toxicology [52] • ATSDR - Case Studies in Environmental Medicine: Arsenic Toxicity [53] • Contaminant Focus: Arsenic [54] by the EPA. • Environmental Health Criteria for Arsenic and Arsenic Compounds, 2001 [55] by the WHO. • Evaluation of the carcinogenicity of arsenic and arsenic compounds [56] by the IARC. • National Institute for Occupational Safety and Health - Arsenic Page [57] • National Pollutant Inventory - Arsenic [58] • origen.net – CCA wood and arsenic: toxicological effects of arsenic [59] • WebElements.com – Arsenic [60] pnb:کنیسرآ
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Article Sources and Contributors
Article Sources and Contributors Arsenic Source: http://en.wikipedia.org/w/index.php?oldid=308216099 Contributors: *drew, 123123john, 21655, 25or6to4, 2D, ACSE, Acolorpink1, Adashiel, Addshore, Ahoerstemeier, Alansohn, Alex earlier account, Alguienboga, Alibb, Alvis, Amble, Amorelli, Ams80, Amsibert, Andre Engels, Andres, Andrew Rodland, Anetode, Angrynight, Anonymous101, Antandrus, Anwar saadat, Arakunem, Aramgutang, Arbitrarily0, Arcadian, Archimerged, ArielGold, Astatine-210, Atokoy, Aussie Alchemist, Austiwang, Avanu, Avono, Awesome91, AxelBoldt, Axiosaurus, Axl, Az1568, Bantman, BarretBonden, Bdodo1992, Bedrupsbaneman, Beetstra, Beland, Belizefan, BenFrantzDale, Benbest, Bendzh, Benwildeboer, BiT, Bigheadedboy, Bikeable, BlueEarth, Bobo192, Bobépierre, Bogey97, Bongwarrior, Booksworm, BoomerAB, Borislav.dopudja, Brandmeister, Brim, Bryan Derksen, Bubba hotep, Bubbachuck, Burntsauce, C3o, CHADMEISTER, CYD, Cacycle, Camembert, Carnildo, Cdang, Ceyockey, Cflm001, CharlotteWebb, Chasingsol, ChemNerd, Chinaexports, Christopher Parham, Chriswaterguy, CiTrusD, Ck lostsword, Closedmouth, Cmdrjameson, Cmichael, Cnmetaltrade, Conversion script, Cpaton, Crestville, Cryptic C62, Cxz111, Cyde, Cyrusuryc, DJBullfish, DVD R W, DanielCD, Darrien, DaveGorman, David Latapie, David Pierce, David spector, Dawn Bard, Dcandeto, Dead3y3, Deli nk, Delirium, Delta G, Djr5353, Dlae, Doc Tropics, Donarreiskoffer, Doniago, Doonhamer, Doulos Christos, Dr. Dunglison, Dreadstar, Drini, Duderusia, Dy yol, ESkog, Ed Poor, Edgar181, Edward, Edwy, Egmonster, El C, Elassint, Eldin raigmore, Elwood j blues, Emperorbma, Epbr123, Epolk, Erik Zachte, Euchiasmus, Evice, Explicit, Fallatio, Faradayplank, Fconaway, Femto, Firefoxman, Floria L, Frank Lofaro Jr., Fredbauder, Freddyd945, Garrettett, Gcsuchemistry, Geht, Gelzo, Giftlite, Gjd001, Gmaxwell, Gogo Dodo, Grassfire, Greatestqueen, Grendelkhan, Gsayles, Gscshoyru, Guanaco, Gumba gumba, Gurch, H1bhaska, Hak-kâ-ngìn, Harish2k1vet, Hburg, Hellbus, Heron, HexaChord, Hippietrail, HokieRNB, Hu12, Hydrogen Iodide, IW.HG, Icairns, Ideyal, Iepeulas, ImperatorExercitus, Inkypaws, Insanity Incarnate, Inspector08, Ioeth, Iridescent, Isis, J.delanoy, J7890bigman, J8079s, JFreeman, JNW, JRPG, Jac16888, Jackcsk, Jagged 85, James500, Jamesontai, Jan.Smolik, Jaraalbe, Jason ost, JavierMC, Jeffrey O. Gustafson, Jermantowicz, Jimjamjak, Jimokay, Jkl, Joanjoc, Johann Wolfgang, John, JohnCD, Johnbibby, Jonny-mt, Jordan042, Jose77, Joshschr, Jossi, Joyous!, Karlhahn, Katieh5584, Keegan, Keilana, KeithH, Kichigoro12, Killiondude, Koalaman520, Kooo, Kpjas, Kralizec!, Krich, Ktsquare, Kubigula, Kukini, Kuru, Kurykh, Kwamikagami, Lee Daniel Crocker, Lightdarkness, LilHelpa, Link2joon, Logicman1966, LuigiManiac, Luk, Lychosis, MER-C, MFago, Magepure892, Mai-tai-guy, Malcolm Farmer, Manette, Mani1, Marcelo.84, Marcelo1229, Marlith, Marnanel, Masoninman, Master Jay, Matt.T, MattieTK, Mav, Melis610, Mgimpel, MightyWarrior, Minesweeper, Mr Bungle, Mr Stephen, Mr. Lefty, MrMunky, Mschel, Mschiffler, MuVo100, Mufka, Natalie Erin, Nburden, NellieBly, Nergaal, Neverquick, NewEnglandYankee, Nibuod, Nick Y., Nick88, Nickj, Nihiltres, Nitchell, Novaprospekt, Number 57, Oajsn, Oliver202, Oneslowlx, Onevalefan, Onions and liver, Opelio, Ossmann, Ottawa4ever, OwenX, Oxymoron83, PDH, PP Jewel, PStatic, Pacific66, Panoptical, PatVanHove, Pedometer+1, Pfahlstrom, Physchim62, Phædrus, Piano non troppo, PierreAbbat, Pishogue, Plasmic Physics, Poolkris, Poor Yorick, Pras, Pstanton, Pstudier, Qxz, RG2, RTC, RainbowOfLight, Raven4x4x, Redux, Remember, Requiems, Reuben, RexNL, Reyk, Rgfjdsf, Riana, Rich Farmbrough, Rifleman 82, Rizzardi, Rjwilmsi, Rmhermen, Rmrfstar, Rob Hooft, Robert McClenon, Roberta F., Romanm, Roscelese, RoyBoy, Rtyq2, Rummmy, SJFriedl, Sander123, Saperaud, Sapphirine, Sbharris, Schneelocke, Sciurinæ, Sehsuan, Sengkang, SexyBern, Shaddack, Shanel, Shizane, Silentlight, Sionus, Sir-Restriction, Sjö, Skatebiker, Sl, Sleigh, Smokizzy, Squids and Chips, Squirepants101, StaticGull, StaticVision, Stemonitis, StephanieM, Stephenb, Steve Crossin, Stevenfruitsmaak, Stifynsemons, Stone, Sunborn, Taivo, Tassedethe, Tempodivalse, Terence, Tetracube, The Minister of War, TheKMan, TheSuperBrain2, Themerejoy, Theseeker4, Thricecube, Thue, Tiddly Tom, Tide rolls, Tim Starling, Tiptoety, Tisdalepardi, Titoxd, Tizzybag, Tkynerd, Tom harrison, Tombadog, Tone, Topdeck, Trevor MacInnis, Triforce of Power, Ttsalo, Turmeric3, Ucanlookitup, Until It Sleeps, Versus22, Victor falk, Vildricianus, Viriditas, Vogon77, Voyagerfan5761, Vsmith, Vuong Ngan Ha, WJBscribe, Wancheseblondie, Warut, Watch37264, Wfructose, Wiki Raja, Wikieditor06, Wolfrock, Yath, Yiplop stick stop, YixilTesiphon, Yyy, Zedla, Zigger, Zsinj, Äpple, Александър, 895 anonymous edits
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14
Selenium
1
Selenium 34
arsenic ← selenium → bromine
S ↑
Se ↓
Te Periodic Table - Extended Periodic Table
General Name, Symbol, Number
selenium, Se, 34
Element category
nonmetals
Group, Period, Block
16, 4, p
Appearance
gray-black, metallic luster
Standard atomic weight
78.96(3) g·mol
Electron configuration
[Ar] 4s 3d
Electrons per shell
−1
2
10
4
4p
2, 8, 18, 6 Physical properties
Phase
solid
Density (near r.t.)
(gray) 4.81 g·cm
Density (near r.t.)
(alpha) 4.39 g·cm−3
Density (near r.t.)
(vitreous) 4.28 g·cm−3
Liquid density at m.p.
3.99 g·cm−3
Melting point
494 K (221 °C, 430 °F)
Boiling point
958 K (685 °C, 1265 °F)
Critical point
1766 K, 27.2 MPa
Heat of fusion
(gray) 6.69 kJ·mol−1
Heat of vaporization
95.48 kJ·mol−1
Specific heat capacity
(25 °C) 25.363 J·mol−1·K−1
−3
Selenium
2
Vapor pressure P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
500
552
617
704
813
958
Atomic properties Crystal structure
hexagonal [1]
Oxidation states
6, 4, 2, 1, -2 (strongly acidic oxide)
Electronegativity
2.55 (Pauling scale)
Ionization energies (more)
1st: 941.0 kJ·mol−1 −1
2nd: 2045 kJ·mol
−1
3rd: 2973.7 kJ·mol Atomic radius
120 pm
Covalent radius
120±4 pm
Van der Waals radius
190 pm Miscellaneous [2]
Magnetic ordering
diamagnetic
Thermal conductivity
(300 K) (amorphous) −1 −1 0.519 W·m ·K
Thermal expansion
(25 °C) (amorphous) −1 −1 37 µm·m ·K
Speed of sound (thin rod)
(20 °C) 3350 m/s
Young's modulus
10 GPa
Shear modulus
3.7 GPa
Bulk modulus
8.3 GPa
Poisson ratio
0.33
Mohs hardness
2.0
Brinell hardness CAS registry number
736 MPa 7782-49-2 Most-stable isotopes
Selenium
3
Main article: Isotopes of selenium iso 72
Se
74
Se
75
Se
NA syn
0.87% syn
half-life 8.4 d
DM ε
-
γ
0.046
119.779 d
9.36%
76
77
7.63%
77
78
23.78%
78
Se
79
Se
syn
80
49.61%
82
8.73%
Se Se
72
As
-
Se is stable with 40 neutron
76
Se
DP
74
ε γ
Se
DE (MeV)
0.264, 0.136, 0.279
75
As
-
Se is stable with 42 neutron Se is stable with 43 neutron Se is stable with 44 neutron 5
2.95×10 y
−
β
0.151
79
2.995
82
Br
80
Se is stable with 46 neutron 20
1.08×10
y
− −
β β
Kr
References
Selenium (pronounced /səˈliːniəm/) is a chemical element with the atomic number 34, represented by the chemical symbol Se, an atomic mass of 78.96. It is a nonmetal, chemically related to sulfur and tellurium, and rarely occurs in its elemental state in nature. Isolated selenium occurs in several different forms, the most stable of which is a dense purplish-gray semi-metal (semiconductor) form that is structurally a trigonal polymer chain. It conducts electricity better in the light than in the dark, and is used in photocells (see allotropes section below). Selenium also exists in many non-conductive forms: a black glass-like allotrope, as well as several red crystalline forms built of eight-membered ring molecules, like its lighter chemical cousin sulfur. Selenium is found in economic quantities in sulfide ores such as pyrite, partially replacing the sulfur in the ore matrix. Minerals that are selenide or selenate compounds are also known, but all are rare. The chief commercial present uses for selenium are in glassmaking and in chemicals and pigments. Electronic uses for selenium, once important, have been supplanted by silicon semiconductor devices. Selenium salts are toxic in large amounts, but trace amounts of the element are necessary for cellular function in most, if not all, animals, forming the active center of the enzymes glutathione peroxidase and thioredoxin reductase (which indirectly reduce certain oxidized molecules in animals and some plants) and three known deiodinase enzymes (which convert one thyroid hormone to another). Selenium requirements in plants differ by species, with some plants apparently requiring none.[3]
Selenium
4
History and global demand Selenium (Greek σελήνη selene meaning "Moon") was discovered in 1817 by Jöns Jakob Berzelius[4] who found the element associated with tellurium (named for the Earth). It was discovered as a byproduct of sulfuric acid production. It came to medical notice later because of its toxicity to humans working in industry. It was also recognized as an important veterinary toxin. In 1954 first hints towards specific biological functions of selenium were discovered in microorganisms. Its essentiality for mammalian life was discovered in 1957. In the 1970s it was shown to be present in two independent sets of enzymes. This was followed by the discovery of selenocysteine in proteins. During the 1980s it was shown that selenocystine was encoded by the codon TGA. The recoding mechanism was worked out first in bacteria and then in mammals. Growth in selenium consumption was historically driven by steady development of new uses, including applications in rubber compounding, steel alloying, and selenium rectifiers. Selenium is also an essential material in the drums of laser printers and copiers. By 1970, selenium in rectifiers had largely been replaced by silicon, but its use as a photoconductor in plain-paper copiers had become its leading application. During the 1980s, the photoconductor application declined (although it was still a large end-use) as more and more copiers using organic photoconductors were produced. Currently, the largest use of selenium worldwide is in glass manufacturing, followed by uses in chemicals and pigments. Electronics use, despite a number of continued applications, continues to decline.[5] In 1996, continuing research showed a positive correlation between selenium supplementation and cancer prevention in humans, but widespread direct application of this important finding would not add significantly to demand owing to the small doses required. In the late 1990s, the use of selenium (usually with bismuth) as an additive to plumbing brasses to meet no-lead environmental standards became important. At present, total world selenium production continues to increase modestly.
Occurrence Selenium occurs naturally in a number of inorganic forms, including selenide, selenate and selenite. In soils, selenium most often occurs in soluble forms such as selenate (analogous to sulfate), which are leached into rivers very easily by runoff.
Native selenium
Selenium has a biological role, and it is found in organic compounds such as dimethyl selenide, selenomethionine, selenocysteine and methylselenocysteine. In these compounds selenium plays a role analogous to that of sulfur.
Selenium is most commonly produced from selenide in many sulfide ores, such as those of copper, silver, or lead. It is obtained as a byproduct of the processing of these ores, from the anode mud of copper refineries and the mud from the lead chambers of sulfuric acid plants. These muds can be processed by a number of means to obtain free selenium. Natural sources of selenium include certain selenium-rich soils, and selenium that has been bioconcentrated by certain plants. Anthropogenic sources of selenium include coal burning and the mining and smelting of sulfide ores.[6]
Selenium
5
See also Selenide minerals.
Production and allotropic forms
From left to right: Black, gray and red selenium
Native selenium is a rare mineral, which does not usually form good crystals, but when it does they are steep rhombohedrons or tiny acicular (hair-like) crystals. [7] Most elemental selenium comes as a byproduct of copper refining, or the production of sulfuric acid.[8] [9] Isolation of selenium is often complicated by the presence of other compounds and elements. Commonly, production begins by oxidation with sodium carbonate to produce selenium dioxide. The selenium dioxide is then mixed with water producing selenous acid. The selenous acid is finally bubbled with sulfur dioxide producing elemental red amorphous selenium.
Selenium produced in chemical reactions invariably appears as the amorphous red form—an insoluble, brick-red powder. When this form is rapidly melted, it forms the black, vitreous form which is usually sold industrially as beads. The most thermodynamically Structure of trigonal selenium stable and dense form of selenium is the electrically conductive gray (trigonal) form, which is composed of long helical chains of selenium atoms (see figure).[10] The conductivity of this form is notably light sensitive. Selenium also exists in three different deep-red crystalline monoclinic forms, which are composed of Se8 molecules, similar to many allotropes of sulfur.[11]
Isotopes Selenium has six naturally occurring isotopes, five of which are stable: 74Se, 76Se, 77Se, 78 Se, and 80Se. The last three also occur as fission products, along with 79Se which has a half-life of 295,000 years. The final naturally occurring isotope, 82Se, has a very long half-life (~1020 yr, decaying via double beta decay to 82Kr) and which for practical purposes can be considered to be stable. 23 other unstable isotopes have been characterized. See also Selenium-79 for more information on recent changes in the half-life of this fission product, important for the dose calculations performed in the frame of the geological disposal of long-lived radioactive waste.
Health effects and nutrition Although it is toxic in large doses, selenium is an essential micronutrient for animals. In plants, it occurs as a bystander mineral, sometimes in toxic proportions in forage (some plants may accumulate selenium as a defense against being eaten by animals, but other plants such as locoweed require selenium, and their growth indicates the presence of selenium in soil).[3] It is a component of the unusual amino acids selenocysteine and selenomethionine. In humans, selenium is a trace element nutrient which functions as
Selenium cofactor for reduction of antioxidant enzymes such as glutathione peroxidases and certain forms of thioredoxin reductase found in animals and some plants (this enzyme occurs in all living organisms, but not all forms of it in plants require selenium). Glutathione peroxidase (GSH-Px) catalyzes certain reactions which remove reactive oxygen species such as peroxide: 2 GSH+ H2O2---------GSH-Px → GSSG + 2 H2O Selenium also plays a role in the functioning of the thyroid gland by participating as a cofactor for the three known thyroid hormone deiodinases.[12] Dietary selenium comes from nuts, cereals, meat, fish, and eggs. Brazil nuts are the richest ordinary dietary source (though this is soil-dependent, since the Brazil nut does not require high levels of the element for its own needs). High levels are found in kidney, tuna, crab and lobster, in that order.[13] [14]
Selenium indicator plants Certain species of plants are considered indicators of high selenium content of the soil, since they require high levels of selenium in order to thrive. The main selenium indicator plants are Astragalus species (including some locoweeds), prince's plume (Stanleya sp.), woody asters (Xylorhiza sp.), and false goldenweed (Oonopsis sp.)[15]
Toxicity Although selenium is an essential trace element, it is toxic if taken in excess. Exceeding the Tolerable Upper Intake Level of 400 micrograms per day can lead to selenosis.[16] This 400 microgram Tolerable Upper Intake Level is primarily based on a 1986 study of five Chinese patients who exhibited overt signs of selenosis and a follow up study on the same five people in 1992.[17] The 1992 study actually found the maximum safe dietary Se intake to be approximately 800 micrograms per day (15 micrograms per kilogram body weight), but suggested 400 micrograms per day to not only avoid toxicity, but also to avoid creating an imbalance of nutrients in the diet and to account for data from other countries.[18] The Chinese people who suffered from selenium toxicity ingested selenium by eating corn grown in extremely selenium-rich stony coal (carbonaceous shale). This coal was shown to have selenium content as high as 9.1%, the highest concentration in coal ever recorded in literature.[19] Symptoms of selenosis include a garlic odor on the breath, gastrointestinal disorders, hair loss, sloughing of nails, fatigue, irritability, and neurological damage. Extreme cases of selenosis can result in cirrhosis of the liver, pulmonary edema and death.[20] Elemental selenium and most metallic selenides have relatively low toxicities because of their low bioavailability. By contrast, selenates and selenites are very toxic, having an oxidant mode of action similar to that of arsenic trioxide. The chronic toxic dose of selenite for human beings is about 2400 to 3000 micrograms of selenium per day for a long time.[21] Hydrogen selenide is an extremely toxic, corrosive gas.[22] Selenium also occurs in organic compounds such as dimethyl selenide, selenomethionine, selenocysteine and methylselenocysteine, all of which have high bioavailability and are toxic in large doses. Nano-size selenium has equal efficacy, but much lower toxicity.[23] [24] [25] [26] [27] [28] [29] On April 19, 2009, twenty-one polo ponies began to die shortly before a match in the United States Polo Open. Three days later, a private pharmacy released a statement that the horses had received an incorrect dose of one of the ingredients used in a vitamin compound
6
Selenium with which the horses had been injected. Such vitamin injections are common to promote recovery after a match. The pharmacy did not initially release the name of the specific ingredient due to ongoing law-enforcement and other investigations. Dr David Barber, an associate professor in the Center for Environmental and Human Toxicology at the University of Florida began conducting analysis of inorganic compounds of the vitamin supplement and discovered that selenium concentrations were ten to fifteen times higher than normal in the horses' blood samples and 15 to 20 times higher than normal in their liver samples. A pharmacy spokesperson later confirmed that selenium was the ingredient in question.[30] Selenium poisoning of water systems may result whenever new agricultural runoff courses through normally dry undeveloped lands. This process leaches natural soluble selenium compounds (such as selenates) into the water, which may then be concentrated in new "wetlands" as the water evaporates. High selenium levels produced in this fashion have been found to have caused certain congenital disorders in wetland birds.[31]
Deficiency Selenium deficiency is relatively rare in healthy, well-nourished individuals. It can occur in patients with severely compromised intestinal function, those undergoing total parenteral nutrition, and also[32] on advanced-aged people (over 90). Alternatively, people dependent on food grown from selenium-deficient soil are also at risk. Interestingly, although New Zealand has low levels of selenium in its soil, adverse health effects have not been detected.[33]
Controversial health effects Cancer Several studies have suggested a possible link between cancer and selenium deficiency,[34] [35] [36] [37] One study, known as the NPC, was conducted to test the effect of selenium supplementation on the recurrence of skin cancers on selenium-deficient men. It did not demonstrate a reduced rate of recurrence of skin cancers, but did show a reduced occurrence of total cancers, although without a statistically significant change in overall mortality.[38] The preventative effect observed in the NPC was greatest in those with the lowest baseline selenium levels.[39] In 2009 the 5.5 year SELECT study reported that selenium and vitamin E supplementation, both alone and together, did not significantly reduce the incidence of prostate cancer in 35,000 men who "generally were replete in selenium at baseline".[39] The SELECT trial found that vitamin E did not reduce prostate cancer as it had in the Alpha-Tocopherol, Beta Carotene (ATBC) study, but the ATBC had a large percentage of smokers while the SELECT trial did not.[39] Dietary selenium prevents chemically induced carcinogenesis in many rodent studies.[40] It has been proposed that selenium may help prevent cancer by acting as an antioxidant or by enhancing immune activity. Not all studies agree on the cancer-fighting effects of selenium. One study of naturally occurring levels of selenium in over 60,000 participants did not show a significant correlation between those levels and cancer.[41] The SU.VI.MAX study[42] concluded that low-dose supplementation (with 120 mg of ascorbic acid, 30 mg of vitamin E, 6 mg of beta carotene, 100 µg of selenium, and 20 mg of zinc) resulted in a 30% reduction in
7
Selenium the incidence of cancer and a 37% reduction in all-cause mortality in males, but did not get a significant result for females.[43] However, there is evidence that selenium can help chemotherapy treatment by enhancing the efficacy of the treatment, reducing the toxicity of chemotherapeutic drugs, and preventing the body's resistance to the [44] drugs. Studies of cancer cells in vitro showed that chemotherapeutic drugs, such as Taxol and Adriamycin, were more toxic to strains of cancer cells grown in culture when selenium was added.[45] [46] In March 2009, a study from the Department of Cancer Biology at the University of Texas M. D. Anderson Cancer Center reports that Vitamin E (400 IU) and selenium (200 micrograms) supplements affect gene expression and can act as a tumor suppressor.[47] Eric Klein, MD from the Glickman Urological and Kidney Institute in Ohio said the new study “lend credence to the previous evidence that selenium and vitamin E might be active as cancer preventatives”.[48] In an attempt to rationalise the differences between epidemiological and in vitro studies and randomised trials like SELECT, Klein said that randomized controlled trials “do not always validate what we believe biology indicates and that our model systems are imperfect measures of clinical outcomes in the real world”.[48] HIV/AIDS Some research has indicated a geographical link between regions of selenium-deficient soils and peak incidences of HIV/AIDS infection. For example, much of sub-Saharan Africa is low in selenium. However, Senegal is not, and also has a significantly lower level of AIDS infection than the rest of the continent. AIDS appears to involve a slow and progressive decline in levels of selenium in the body. Whether this decline in selenium levels is a direct result of the replication of HIV[49] or related more generally to the overall malabsorption of nutrients by AIDS patients remains debated. Low selenium levels in AIDS patients have been directly correlated with decreased immune cell count and increased disease progression and risk of death.[50] Selenium normally acts as an antioxidant, so low levels of it may increase oxidative stress on the immune system leading to more rapid decline of the immune system. Others have argued that T-cell associated genes encode selenoproteins similar to human glutathione peroxidase. Depleted selenium levels in turn lead to a decline in CD4 helper T-cells, further weakening the immune system.[51] Regardless of the cause of depleted selenium levels in AIDS patients, studies have shown that selenium deficiency does strongly correlate with the progression of the disease and the risk of death.[52] [53] [54] Tuberculosis Some research has suggested that selenium supplementation, along with other nutrients, can help prevent the recurrence of tuberculosis.[55] Diabetes A well-controlled study showed that selenium intake is positively correlated with the risk of developing type II diabetes. Because high serum selenium levels are positively associated with the prevalence of diabetes, and because selenium deficiency is rare, supplementation is not recommended in well-nourished populations such as the U.S.[56] Mercury
8
Selenium Experimental findings have demonstrated a protective effect of selenium on methylmercury toxicity, but epidemiological studies have been inconclusive in linking selenium to protection against the adverse effects of methylmercury.[57]
Non-biologic applications Chemistry Selenium is a catalyst in many chemical reactions and is widely used in various industrial and laboratory syntheses, especially Organoselenium chemistry. It is also widely used in structure determination of proteins and nucleic acids by X-ray crystallography (incorporation of one or more Se atoms helps with MAD and SAD phasing.) Manufacturing and materials use The largest use of selenium worldwide is in glass and ceramic manufacturing, where it is used to give a red color to glasses, enamels and glazes as well as to remove color from glass by counteracting the green tint imparted by ferrous impurities. Selenium is used with bismuth in brasses to replace more toxic lead. It is also used to improve abrasion resistance in vulcanized rubbers. Electronics Because of its photovoltaic and photoconductive properties, selenium is used in photocopying, photocells, light meters and solar cells. It was once widely used in rectifiers. These uses have mostly been replaced by silicon-based devices, or are in the process of being replaced. The most notable exception is in power DC surge protection, where the superior energy capabilities of selenium suppressors make them more desirable than metal oxide varistors. Sheets of amorphous selenium convert x-ray images to patterns of charge in xeroradiography and in solid-state, flat-panel x-ray cameras. Photography Selenium is used in the toning of photographic prints, and it is sold as a toner by numerous photographic manufacturers including Kodak and Fotospeed. Its use intensifies and extends the tonal range of black and white photographic images as well as improving the permanence of prints.
Biologic applications Medical use The substance loosely called selenium sulfide, SeS2, actually selenium disulfide or selenium (IV) sulfide, is the active ingredient in some dandruff shampoos.[58] The effect of the active ingredient is to kill the scalp fungus Malassezia which causes shedding of dry skin fragments. The ingredient is also used in body lotions to treat Tinea versicolor due to infection by a different species of Malassezia fungus.[59] Nutrition Selenium is used widely in vitamin preparations and other dietary supplements, in small doses (typically 50 to 200 micrograms per day for adult humans). Some livestock feeds are fortified with selenium as well.
9
Selenium
Evolution in biology Over three billion years ago, blue-green algae were the most primitive oxygenic photosynthetic organisms and are ancestors of multicellular eukaryotic algae.[60] Algae that contain the highest amount of antioxidant selenium, iodide, and peroxidase enzymes were the first living cells to produce poisonous oxygen in the atmosphere. Venturi et al.[60] [61] suggested that algal cells required a protective antioxidant action, in which selenium and iodides, through peroxidase enzymes, have had this specific role. Selenium, which acts synergistically with iodine,[62] is a primitive mineral antioxidant, greatly present in the sea and prokaryotic cells, where it is an essential component of the family of glutathione peroxidase antioxidant enzymes (GSH-Px). In fact, seaweeds accumulate high quantity of selenium and iodine.[60] In 2008, Küpper et al.,[63] showed that iodide also scavenges reactive oxygen species (ROS) in algae, and that its biological role is that of an inorganic antioxidant, the first to be described in a living system, active also in an in vitro assay with the blood cells of today’s humans." From about three billion years ago, prokaryotic selenoprotein families drive selenocysteine evolution. Selenium is incorporated into several prokaryotic selenoprotein families in bacteria, archaea and eukaryotes as selenocysteine,[64] where selenoprotein peroxiredoxins protect bacterial and eukaryotic cells against oxidative damage. Selenoprotein families of GSH-Px and deiodinase of eukaryotic cells seem to have a bacterial phylogenetic origin. The selenocysteine-containing form occurred in green algae, diatoms, sea urchin, fish and chicken, too. One family of selenium-containing molecules as glutathione peroxidases repairs damaged cell membranes, while another (glutathione S-transferases) repairs damaged DNA and prevents mutations.[65] When about 500 Mya, plants and animals began to transfer from the sea to rivers and land, the environmental deficiency of marine mineral antioxidants (as selenium, iodine, etc.) was a challenge to the evolution of terrestrial life.[60] Trace elements involved in GSH-Px and superoxide dismutase enzymes activities, i.e. selenium, vanadium, magnesium, copper and zinc, may have been lacking in some terrestrial mineral-deficient areas.[64] Marine organisms apparently retained and sometimes expanded their seleno-proteomes, whereas the seleno-proteomes of some terrestrial organisms were reduced or completely lost. These findings suggest that, with the exception of vertebrates, aquatic life supports selenium utilization, whereas terrestrial habitats lead to reduced use of this trace element.[66] Marine fishes and vertebrate thyroid glands have the highest concentration of selenium and iodine. From about 500 Mya, freshwater and terrestrial plants slowly optimized the production of “new” endogenous antioxidants such as ascorbic acid (Vitamin C), polyphenols, flavonoids, tocopherols, etc. A few of these appeared more recently, in the last 50-200 million years, in fruits and flowers of angiosperm plants. In fact, the angiosperms (the dominant type of plant today) and most of their antioxidant pigments evolved during the late Jurassic period. The deiodinase isoenzymes constituted the second family of eukaryotic selenoproteins with identified enzyme function. Deiodinases are able to extract electrons from iodides, and iodides from iodothyronines; so, they are involved in thyroid-hormone regulation, participating in the protection of thyrocytes from damage by H2O2 produced for thyroid-hormone biosynthesis.[60] [61] About 200 Mya, new selenoproteins were developed as mammalian GSH-Px enzymes.[67] [68] [69] [70]
10
Selenium
11
Compounds • • • • • • • • • •
Copper indium gallium selenide Cu(Ga,In)Se2 Mercury selenide (HgSe) Hydrogen selenide (H2Se) Lead selenide (PbSe) Selenium dioxide (SeO2) Selenic acid (H2SeO4) Selenous acid (H2SeO3) Selenium sulfides: Se4S4, SeS2, Se2S6 Sodium selenite (Na2SeO3) Zinc selenide (ZnSe)
Selenium occurs in the 0,+2,+4,+6 and -2 valence states. See also Selenium compounds and organoselenium chemistry.
External links • WebElements.com - Selenium
[71]
• National Institutes of Health page on Selenium • ATSDR - Toxicological Profile: Selenium [73] • Peter van der Krogt elements site [74]
[72]
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http:/ / www. webelements. com/ webelements/ elements/ text/ Se/ index. html http:/ / ods. od. nih. gov/ factsheets/ selenium. asp http:/ / www. atsdr. cdc. gov/ toxprofiles/ tp92. html http:/ / elements. vanderkrogt. net/ elem/ se. html
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15
Bromine
1
Bromine 35
selenium ← bromine → krypton
Cl ↑
Br ↓
I Periodic Table - Extended Periodic Table
General Name, Symbol, Number
bromine, Br, 35
Element category
halogens
Group, Period, Block
17, 4, p
Appearance
gas/liquid: red-brown solid: metallic luster
Standard atomic weight
79.904(1) g·mol
Electron configuration
[Ar] 4s 3d
−1
2
Electrons per shell
10
5
4p
2, 8, 18, 7 Physical properties
Phase
liquid
Density (near r.t.)
(Br , liquid) 3.1028 g·cm−3
Melting point
265.8 K (-7.2 °C, 19 °F)
Boiling point
332.0 K (58.8 °C, 137.8 °F)
Critical point
588 K, 10.34 MPa
Heat of fusion
(Br ) 10.571 kJ·mol−1
2
2
Heat of vaporization
(Br ) 29.96 kJ·mol−1 2
Specific heat capacity
(25 °C) (Br ) 2 75.69 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
185
201
220
244
276
332
Atomic properties
Bromine
2
Crystal structure
orthorhombic
Oxidation states
7, 5, 4, 3, 1, -1 (strongly acidic oxide)
Electronegativity
2.96 (Pauling scale)
Ionization energies (more)
1st: 1139.9 kJ·mol−1 2nd: 2103 kJ·mol−1 3rd: 3470 kJ·mol−1
Atomic radius
120 pm
Covalent radius
120±3 pm
Van der Waals radius
185 pm Miscellaneous [1]
Magnetic ordering
diamagnetic
Electrical resistivity
(20 °C) 7.8×10
Thermal conductivity
(300 K) 0.122 W·m
Speed of sound
(20 °C) ? 206 m/s
10
Ω·m −1
CAS registry number
−1
·K
7726-95-6 Most-stable isotopes
Main article: Isotopes of bromine iso
NA
half-life
79
50.69%
79
81
49.31%
81
Br Br
DM
DE (MeV)
DP
Br is stable with 44 neutron Br is stable with 46 neutron References
Bromine (pronounced /ˈbroʊmiːn/ or English pronunciation: /ˈbroʊmɨn/), Greek: βρῶμος, brómos, meaning "stench (of he-goats)" [2] ), is a chemical element with the symbol Br and atomic number 35. A halogen element, bromine is a reddish-brown volatile liquid at standard room temperature that is intermediate in reactivity between chlorine and iodine. Bromine vapours are corrosive and toxic. Approximately 556,000 metric tons were produced in 2007. [3] The main applications for bromine are in fire retardants and fine chemicals.
Bromine
3
History Bromine was discovered independently by two chemists Antoine Balard[4] and Carl Jacob Löwig[5] in 1825 and 1826.[6]
Illustrative and secure bromine sample for teaching
Balard found bromide salts in the ash of sea weed from the salt marshes of Montpellier in 1826. The sea weed was used to produce iodine, but also contained bromine. Balard distilled the bromine from a solution of seaweed ash saturated with chlorine. The properties of the resulting substance resembled that of an intermediate of chlorine and iodine, with those results he tried to prove that the substance was iodine monochloride (ICl), but after failing to do so he was sure that he found a new element and named it muride, derived from the Latin word muria for brine.[4]
Carl Jacob Löwig isolated bromine from mineral water spring from his home town Bad Kreuznach in 1825. Löwig used a solution of the mineral salt saturated with chlorine and extracted the bromine with diethylether. After evaporation of the ether a brown liquid remained. With this liquid as a sample for his work he applied for a position in the laboratory of Leopold Gmelin in Heidelberg. The publication of the results was delayed and Balard published his results first.[5] After the French chemists Louis Nicolas Vauquelin, Louis Jacques Thénard, and Joseph-Louis Gay-Lussac approved the experiments of the young pharmacist Balard, the results where presented at a lecture of the Académie des Sciences and published in Annales de Chimie et Physique[7] . In his publication Balard states that he changed the name from muride to brôme on the proposal of M. Anglada. Other sources claim that the French chemist and physicist Joseph-Louis Gay-Lussac suggested the name brôme due to the characteristic smell of the vapors.[8] Bromine was not produced in large quantities until 1860. The first commercial use, besides some minor medical applications, was the use of bromine for the daguerreotype. In 1840 it was discovered that bromine had some advantages over the previous used iodine vapour to create the light sensitive silver halide layer used for daguerreotypy.[9] Potassium bromide and sodium bromide were used as anticonvulsants and sedatives in the late 19th and early 20th centuries, until they were gradually superseded by chloral hydrate and then the barbiturates.[10]
Characteristics Bromine is the only liquid nonmetallic element at room temperature, and one of only two elements on the periodic table that are liquid at room temperature. The melting point of bromine is −7.2 °C and has the boiling point 58.8 °C. The pure chemical element has the physical form of a diatomic molecule, Br2. It is a dense, mobile, reddish-brown liquid, that evaporates easily at standard temperature and pressures to give a red vapor (its color resembles nitrogen dioxide) that has a strong disagreeable odor resembling that of chlorine. Bromine is a halogen, and is less reactive than chlorine and more reactive than
Bromine
4
iodine. Bromine is slightly soluble in water, and highly soluble in carbon disulfide, aliphatic alcohols (such as methanol), and acetic acid. It bonds easily with many elements and has a strong bleaching action. Bromine, like chlorine, is also used in maintenance of swimming pools. Certain bromine-related compounds have been evaluated to have an ozone depletion potential or bioaccumulate in living organisms. As a result many industrial bromine compounds are no longer manufactured, are being restricted, or scheduled for phasing out. The Montreal Protocol mentions several organobromine compounds for this phase out. Bromine is a powerful oxidizing agent. It reacts vigorously with metals, especially in the presence of water, as well as most organic compounds, especially upon illumination.
Isotopes Bromine has 2 stable isotopes: 79Br (50.69 %) and 81 Br (49.31%). At least another 23 radioisotopes are known to exist.[11] Many of the bromine isotopes are fission products. Several of the heavier bromine isotopes from fission are delayed neutron emitters. All of the radioactive bromine isotopes are relatively short lived. The longest half life is the neutron deficient 77Br at 2.376 days. The longest half life on the neutron rich side is 82Br at 1.471 days. A number of the bromine isotopes exhibit metastable isomers. Stable 79Br exhibits a radioactive isomer, with a half life of 4.86 seconds. It decays by isomeric transition to the stable ground state.[12]
Occurrence and production See also Halide minerals. The diatomic element Br
2
does not occur naturally.
Instead, bromine exists exclusively as bromide salts in diffuse amounts in crustal rock. Due to leaching, bromide salts have accumulated in sea water (65 ppm),[13] but at a lower concentration than chloride. Bromine may be economically recovered from bromide-rich brine wells and from the Dead Sea waters (up to 50000 ppm).[14] [15] World bromine production trend
Bromine
5
Approximately 556,000 metric tons (worth around US$2.5 billion) of bromine are produced per year (2007) worldwide with the United States, Israel, and China being the primary producers.[16] [17] [18] Bromine production has increased sixfold since the 1960s. The largest bromine reserve in the United States is located in Columbia and Union County, Arkansas, U.S.[19] China's bromine reserves are located in the Shandong Province and Israel's bromine reserves are contained in the waters of the Dead Sea. The bromide-rich brines are treated with chlorine gas, flushing through with air. In this treatment, bromide anions are oxidized to bromine by the chlorine gas. 2 Br− + Cl2 → 2 Cl− + Br2
View of salt evaporation pans on the Dead Sea, where Jordan (right) and Israel (left) produce salt and bromine
Because of its commercial availability and long shelf-life, bromine is not typically prepared. Small amounts of bromine can however be generated through the reaction of solid sodium bromide with concentrated sulfuric acid (H2SO4). The first stage is formation of hydrogen bromide (HBr), which is a gas, but under the reaction conditions some of the HBr is oxidized further by the sulfuric acid to form bromine (Br2) and sulfur dioxide (SO2). NaBr (s) + H2SO4 (aq) → HBr (aq) + NaHSO4 (aq) 2 HBr (aq) + H2SO4 (aq) → Br2 (g) + SO2 (g) + 2 H2O (l) Similar alternatives, such as the use of dilute hydrochloric acid with sodium hypochlorite, are also available. The most important thing is that the anion of the acid (in the above examples, sulfate and chloride, respectively) be more electronegative than bromine, allowing the substitution reaction to occur. Reaction between a strong oxidizing agent such as potassium permanganate on bromide ions in the presence of an acid also gives bromine. An acidic solution of bromate ions and bromide ions will also disproportionate slowly to give bromine. Bromine is only slightly soluble in water. But the solubility can be increased by the presence of bromide ions. However, concentrated solutions of bromine are rarely prepared in the lab as they will continually give off toxic red-brown bromine gas due to its very high vapour pressure. Sodium thiosulphate is an excellent reagent for getting rid of bromine completely including the stains and odour.
Bromine
6
Compounds Organic chemistry Organic compounds are brominated by either addition or substitution reactions. Bromine undergoes electrophilic addition to the double-bonds of alkenes, via a cyclic bromonium intermediate. In non-aqueous solvents such as carbon disulfide, this affords the di-bromo product. For example, reaction with ethylene will produce N-Bromosuccinimide 1,2-dibromoethane. Bromine also undergoes electrophilic addition to phenols and anilines. When used as bromine water, a small amount of the corresponding bromohydrin is formed as well as the dibromo compound. So reliable is the reactivity of bromine that bromine water is employed as a reagent to test for the presence of alkenes, phenols, and anilines. Like the other halogens, bromine participates in free radical reactions. For example hydrocarbons are brominated upon treatment with bromine in the presence of light. Bromine, sometimes with a catalytic amount of phosphorus, easily brominates carboxylic acids at the α-position. This method, the Hell-Volhard-Zelinsky reaction, is the basis of the commercial route to bromoacetic acid. N-Bromosuccinimide is commonly used as a substitute for elemental bromine, being easier to handle, and reacting more mildly and thus more selectively. Organic bromides are often preferable relative to the less reactive chlorides and more expensive iodide-containing reagents. Thus, Grignard and organolithium compound are most often generated from the corresponding bromides. Inorganic chemistry Oxidation states of bromine -1
HBr
+1
BrCl
+3
BrF3
+5
BrF5
+5
BrO−3
+7
BrO−4
Bromine is an oxidizer, and it will oxidize iodide ions to iodine, being itself reduced to bromide: Br2 + 2 I− → 2 Br− + I2 Bromine will also oxidize metals and metalloids to the corresponding bromides. Anhydrous bromine is less reactive toward many metals than hydrated bromine, however. Dry bromine reacts vigorously with aluminium, titanium, mercury as well as alkaline earths and alkali metals. If bromine is dissolved in hydroxide containing water not only bromide (Br−) is formed, but also the hypobromite (OBr−). This hypobromite is responsible for the bleaching abilities of bromide solutions. In warm solutions the disproportion reaction of the hypobromite is quantitative. The resulting bromate is a strong oxidation reagent and very similar to the chlorate.
Bromine
7 3 OBr− → BrO−3 + 2 Br−
The perbromates are not accessible through electrolysis like the perchlorates, but only by reacting bromate solutions with fluorine. OBr− + H2O + F2 → BrO4 + 2 HF
Applications A wide variety of organobromine compounds are used in industry. Some are prepared from bromine and others are prepared from hydrogen bromide, which is obtained by burning hydrogen in bromine.[3] Illustrative of the addition reaction[20] is the preparation of 1,2-Dibromoethane, the organobromine compound produced in the largest amounts: C2H4 + Br2 → CH2BrCH2Br
Flame retardant Brominated flame retardants represent a commodity of growing importance. If the material burns the flame retardants produce hydrobromic acid which interferes in the radical chain reaction of the oxidation reaction of the fire. The highly reactive hydrogen oxygen and hydroxy radicals react with hydrobromic acid and form less reactive bromine radicals.[21] [22] The bromine Tetrabromobisphenol A containing compounds can be placed in the polymers either during polymerisation if a small amount of brominated monomer is added or the bromine containing compound is added after polymerisation. Tetrabromobisphenol A can be added to produce polyesters or epoxy resins. An epoxy resigns used for printed circuit boards (PCB) are normally made from flame retardant resigns, indicated by the FR in the abbreviation of the products (FR-4 and FR-2. Vinyl bromide can be used in the production of polyethylene, polyvinylchloride or polypropylene. Decabromodiphenyl ether can be added to the final polymers.[23]
Gasoline additive Ethylene bromide was an additive in gasolines containing lead anti-engine knocking agents. It scavenges lead by forming volatile lead bromide, which is exhausted from the engine. This application accounted for 77% of the bromine uses in 1966 in the US. This application has declined since the 1970s due to environmental regulations.[24] Ethylene bromide is also used as a fumigant, but again this application is declining.[18]
Bromine
8
Pesticide Methyl bromide was widely used as pesticide to fumigate soil. The Montreal Protocol on Substances that Deplete the Ozone scheduled the phase out for the ozone depleting chemical until 2005. In 1991, an estimated 35,000 metric tonnes of the chemical were used to control nematodes, fungi, weeds and other soil-borne diseases.[25] [26] Methyl bromide (Bromomethane)
Other Use • The bromides of calcium, sodium, and zinc account for a sizable part of the bromine market. These salts form dense solutions in water that are used as drilling fluids sometimes called clear brine fluids.[18] [27] • Bromine is also used in for the production of brominated vegetable oil, which is used as an emulsifier in many citrus-flavored soft Orange fluoresces of drinks (e.g. Mountain Dew). After the introduction in the 1940s the DNA Ethidium compound was extensively used until the UK and the US limited its bromide intercalate use in the mid 1970s and alternative emulsifers were developed.[28] By 1997 in the US still soft drinks are available containing brominated vegetable oil.[29] • Several dyes, agrichemicals, and pharmaceuticals are organobromine compounds. 1-Bromo-3-chloropropane, 1-bromoethylbenzene, and 1-bromoalkanes are prepared by the antimarkovnikov addition of HBr to alkenes. Ethidium bromide, EtBr, is used as a DNA stain in gel electrophoresis. • High refractive index compounds • Water purification compounds, Disinfectants[18]
Tralomethrin
• Potassium bromide is used in some photographic developers to inhibit the formation of fog (undesired reduction of silver). • Vapor is used as the second step in sensitizing daguerreotype plates to be developed under Mercury (Hg) vapor. Bromine acts as an accelerator to the light sensitivity of the previously iodized plate.
Bromine
9
Biological role Bromine has no known role in human health, but organobromine compounds do occur naturally. Marine organisms are the main source of organobromine compounds. In 1999 over 1600 compounds were identified. The most abundant one is methyl bromide with estimated 56000 metric tons produced by marine algae.[30] The essential oil of the Hawaiian alga Asparagopsis taxiformis consists of 80% methyl bromide.[31] A famous example of a bromine-containing organic compound that has been used by humans for a long time is Tyrian purple.[30] [32] The brominated indigo is produced by a medium-sized Tyrian purple predatory sea snail, the marine gastropod Murex brandaris. It took until 1909 before the organobromine nature of the compound was discovered by Paul [33] Friedländer. Most organobromine compounds in nature arise via the action of vanadium bromoperoxidase.[34]
Safety Elemental bromine is toxic and causes burns. As an oxidizing agent, it is incompatible with most organic and inorganic compounds. Care needs to be taken when transporting bromine; it is commonly carried in steel tanks lined with lead, supported by strong metal frames. When certain ionic compounds containing bromine are mixed with potassium permanganate (KMnO4) and an acidic substance, they will form a pale brown cloud of bromine gas. This gas smells like bleach and is very irritating to the mucus membranes. Upon exposure, one should move to fresh air immediately. If symptoms arise, medical attention is needed.
External links • WebElements.com – Bromine • • • • •
[35]
Theodoregray.com – Bromine [36] USGS Minerals Information: Bromine [37] Bromine Science and Environmental Forum (BSEF) Thermal Conductivity of BROMINE [39] Viscosity of Bromine [40]
[38]
Bromine
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] Gemoll W, Vretska K (1997). Griechisch-Deutsches Schul- und Handwörterbuch ("Greek-German dictionary"), 9th ed.. öbvhpt. ISBN 3-209-00108-1. [3] Jack F. Mills (2002). Bromine: in Ullmann's Encyclopedia of Chemical Technology. Weinheim: Wiley-VCH Verlag. doi: 10.1002/14356007.a04_391 (http:/ / dx. doi. org/ 10. 1002/ 14356007. a04_391). [4] Balard, Antoine (1826). " Memoire of a peculire Substance contained in Sea Water (http:/ / books. google. com/ books?id=A-M4AAAAMAAJ)". Annals of Philosophy: 387– and 411–. . [5] Landolt, Hans Heinrich (1890). " Nekrolog: Carl Löwig (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k907222/ f920. chemindefer)". Berichte der deutschen chemischen Gesellschaft 23: 905. doi: 10.1002/cber.18900230395 (http:/ / dx. doi. org/ 10. 1002/ cber. 18900230395). . [6] Weeks, Mary Elvira (1932). "The discovery of the elements: XVII. The halogen family.". Journal of Chemical Education 9: 1915. [7] Balard, A.J. (1826). Annales de Chimie et Physique 32: 337. [8] Wisniak, Jaime (2004). " Antoine-Jerôme Balard. The discoverer of bromine (http:/ / revistas. mes. edu. cu:9900/ EDUNIV/ 03-Revistas-Cientificas/ Rev. CENIC-Ciencias-Quimicas/ 2004/ 1/ 09204109. pdf)". Revista CENIC Ciencias Químicas 35. . [9] Barger, M. Susan; White, William Blaine (2000). "Technological Practice of Daguerreotypy". The Daguerreotype: Nineteenth-century Technology and Modern Science. JHU Press. pp. 31–35. ISBN 9780801864582. [10] Shorter, Edward (1997). A History of Psychiatry: From the Era of the Asylum to the Age of Prozac. John Wiley and Sons. p. 200. ISBN 9780471245315. [11] GE (1989). Chart of the Nuclides, 14th Edition. Nuclear Energy. [12] Audi, Georges (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [13] Tallmadge, John A.; Butt, John B.; Solomon Herman J. (1964). "Minerals From Sea Salt". Ind. Eng. Chem. 56: 44. doi: 10.1021/ie50655a008 (http:/ / dx. doi. org/ 10. 1021/ ie50655a008). [14] Oumeish, Oumeish Youssef (1996). "Climatotherapy at the Dead Sea in Jordan". Clinics in Dermatology 14: 659. doi: 10.1016/S0738-081X(96)00101-0 (http:/ / dx. doi. org/ 10. 1016/ S0738-081X(96)00101-0). [15] Al-Weshah, Radwan A. (2008). "The water balance of the Dead Sea: an integrated approach". Hydrological Processes 14: 145. doi: 10.1002/(SICI)1099-1085(200001)14:1<145::AID-HYP916>3.0.CO;2-N (http:/ / dx. doi. org/ 10. 1002/ (SICI)1099-1085(200001)14:1<145::AID-HYP916>3. 0. CO;2-N). [16] Emsley, John (2001). "Bromine". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. pp. 69-73. ISBN 0198503407. [17] Lyday, Phyllis A.. " Comodity Report 2007: Bromine (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ bromine/ bromimcs07. pdf)". United States Geological Survey. . Retrieved 2008-09-03. [18] Lyday, Phyllis A.. " Mineral Yearbook 2007: Bromine (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ bromine/ myb1-2006-bromi. pdf)". United States Geological Survey. . Retrieved 2008-09-03. [19] " Bromine:An Important Arkansas Industry (http:/ / www. cals. lib. ar. us/ butlercenter/ lesson_plans/ lesson plans/ Lesson plans-retained/ Bromine. pdf)". . [20] N. A. Khan, F. E. Deatherage, and J. B. Brown (1963). " Stearolic Acid (http:/ / www. orgsyn. org/ orgsyn/ orgsyn/ prepContent. asp?prep=CV4P0851)". Org. Synth.; Coll. Vol. 4: 851. [21] Green, Joseph (1996). "Mechanisms for Flame Retardancy and Smoke suppression -A Review". Journal of Fire Sciences 14: 426. doi: 10.1177/073490419601400602 (http:/ / dx. doi. org/ 10. 1177/ 073490419601400602). [22] Kaspersma, Jelle; Doumena, Cindy; Munrob Sheilaand; Prinsa, Anne-Marie (2002). "Fire retardant mechanism of aliphatic bromine compounds in polystyrene and polypropylene". Polymer Degradation and Stability 77: 325. doi: 10.1016/S0141-3910(02)00067-8 (http:/ / dx. doi. org/ 10. 1016/ S0141-3910(02)00067-8). [23] Weil, Edward D.; Levchik, Sergei (2004). "A Review of Current Flame Retardant Systems for Epoxy Resins". Journal of Fire Sciences 22: 25. doi: 10.1177/0734904104038107 (http:/ / dx. doi. org/ 10. 1177/ 0734904104038107). [24] Alaeea, Mehran; Ariasb, Pedro; Sjödinc, Andreas; Bergman, Åke (2003). "An overview of commercially used brominated flame retardants, their applications, their use patterns in different countries/regions and possible modes of release". Environment International 29: 683. doi: 10.1016/S0160-4120(03)00121-1 (http:/ / dx. doi. org/ 10. 1016/ S0160-4120(03)00121-1). [25] Messenger, Belinda; Braun, Adolf (2000). " Alternatives to Methyl Bromide for the Control of Soil-Borne Diseases and Pests in California (http:/ / www. cdpr. ca. gov/ docs/ emon/ methbrom/ alt-anal/ sept2000. pdf)".
10
Bromine Pest Management Analysis and Planning Program. . Retrieved 2008-11-17. [26] Decanio, Stephen J.; Norman, Catherine S. (2008). "Economics of the "Critical Use" of Methyl bromide under the Montreal Protocol". Contemporary Economic Policy 23: 376. doi: 10.1093/cep/byi028 (http:/ / dx. doi. org/ 10. 1093/ cep/ byi028). [27] Darley, H. C. H.; Gray, George Robert (1988). Composition and Properties of Drilling and Completion Fluids. Gulf Professional Publishing. pp. 61-62. ISBN 9780872011472. [28] Kaufman, Vered R.; Garti, Nissim (1984). "Effect of cloudy agents on the stability and opacity of cloudy emulsions for soft drinks". International Journal of Food Science & Technology 19: 255. doi: 10.1111/j.1365-2621.1984.tb00348.x (http:/ / dx. doi. org/ 10. 1111/ j. 1365-2621. 1984. tb00348. x). [29] Horowitz, B. Zane (1997). "Bromism from Excessive Cola Consumption',Clinical Toxicology". Clinical Toxicology 35: 315. doi: 10.3109/15563659709001219 (http:/ / dx. doi. org/ 10. 3109/ 15563659709001219). [30] Gordon W. Gribble (1999). "The diversity of naturally occurring organobromine compounds". Chemical Society Reviews 28: 335. doi: 10.1039/a900201d (http:/ / dx. doi. org/ 10. 1039/ a900201d). [31] Burreson, B. Jay; Moore, Richard E.; Roller, Peter P. (1976). "Volatile halogen compounds in the alga Asparagopsis taxiformis (Rhodophyta)". Journal of Agricultural snd Food Chemistry 24: 856. doi: 10.1021/jf60206a040 (http:/ / dx. doi. org/ 10. 1021/ jf60206a040). [32] Gordon W. Gribble (1998). "Naturally Occurring Organohalogen Compounds". Acc. Chem. Res. 31: 141. doi: 10.1021/ar9701777 (http:/ / dx. doi. org/ 10. 1021/ ar9701777). [33] Friedländer, P. (1909). "Über den Farbstoff des antiken Purpurs aus murex brandaris". Berichte der deutschen chemischen Gesellschaft 42: 765. doi: 10.1002/cber.190904201122 (http:/ / dx. doi. org/ 10. 1002/ cber. 190904201122). [34] Butler, Alison; Carter-Franklin, Jayme N. (2004). "The role of vanadium bromoperoxidase in the biosynthesis of halogenated marine natural products". Natural Product Reports 21: 180. doi: 10.1039/b302337k (http:/ / dx. doi. org/ 10. 1039/ b302337k). [35] [36] [37] [38] [39] [40]
http:/ / www. webelements. com/ webelements/ elements/ text/ Br/ index. html http:/ / www. theodoregray. com/ PeriodicTable/ Elements/ 035/ index. s7. html http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ bromine http:/ / www. bsef. com/ http:/ / twt. mpei. ac. ru/ MAS/ Worksheets/ HEDH/ 5-5-14-43-54/ Tab-5-5-14-54-BROMINE-Thermal. mcd http:/ / twt. mpei. ac. ru/ MAS/ Worksheets/ HEDH/ 5-5-14-43-54/ Tab-5-5-14-54-BROMINE-Viscosity. mcd
11
Article Sources and Contributors
Article Sources and Contributors Bromine Source: http://en.wikipedia.org/w/index.php?oldid=305255565 Contributors: 2D, A Softer Answer, A. Carty, Ace of Spades IV, Aeusoes1, Ahoerstemeier, Aitias, Ajaxkroon, Akldawgs, Alansohn, Alchemist-hp, Alex43223, Ambuj.Saxena, Amitch, AndonicO, Andres, Antandrus, Archimerged, Asparagusfactory, Aude, B07, Bassbonerocks, Beetstra, Benbest, Bengisuk, Bentu, Blanchardb, BlueEarth, Boat5210, Bobak, Bobo192, Bodnotbod, Bomac, Bonchickawannabowdowlol43434555423, Bonorox, Boud, Brain40, Brandonrush, Brockert, Bryan Derksen, Bunny Angel13, C'est moi, CALR, CWii, CYD, Can't sleep, clown will eat me, Canthusus, CapitalR, Celarnor, ChemNerd, Chrislk02, Christopherk, Chriswaterguy, Chuljin, Colbuckshot, Conversion script, Cryptic C62, DStoykov, DanielCD, Danielsavoiu, DarkHorizon, Darklilac, Darrien, David Latapie, DeadEyeArrow, Dilbatt, DocWatson42, DoubleBlue, Dumelow, Dwmyers, Eastlaw, Eatcacti, Edgar181, Edsanville, El C, Emperorbma, Enviroboy, Epbr123, Erik Zachte, Escape Orbit, Explicit, Femto, Firq, Fl, Flakeloaf, Fortran, Frank Warmerdam, Frankenpuppy, FreeKresge, FreplySpang, Fuzzform, Gary King, Giftlite, Gilliam, Glenn, Greenhorn1, Grendelkhan, Gurch, Gwernol, H Padleckas, Hadal, Hairchrm, HalfShadow, Heartonsleeve, Hellbus, HenryLi, Hermione1980, Holocron, Horkenbubble, Hplommer, Hydrogen Iodide, I'm with gerrit, II MusLiM HyBRiD II, Icairns, Ikiroid, Insanity Incarnate, J.delanoy, J3bus, JForget, Jackol, JamesMLane, Jaraalbe, Jason Quinn, Jdavidb, Jeff G., Jguad1, Jj137, Joanjoc, John, Joshuaali, Jredmond, Juliancolton Alternative, Kakofonous, Karlhahn, Karuna8, Keenan Pepper, Keilana, Keith Edkins, KnowledgeOfSelf, KokomoNYC, Kostisl, Krsloat, Kukini, Kumorifox, Kungfuadam, Kuru, Kurykh, Kwamikagami, Kyle1278, LA2, LarryMorseDCOhio, LeaveSleaves, Lectonar, LeighvsOptimvsMaximvs, Leyo, Liastnir, LilHelpa, Loren.wilton, LuigiManiac, Luk, LukeSurl, Macaddct1984, MacsBug, Macy, Malbi, Materialscientist, Math Champion, Mathx314, Mattlore, Mav, Maxis ftw, Melvinkool.93, Mentisock, Mgimpel, Minesweeper, Mkweise, Mr0t1633, Mrholybrain, Mrsppeed, Mysdaao, NERIC-Security, NawlinWiki, Nergaal, Netalarm, Neurolysis, Nihiltres, Nmnogueira, No1lakersfan, Nornen3, Ottawa4ever, OwenX, Oxymoron83, PastaDruid, Paunaro, PeterJeremy, PeterSymonds, PhageRules1, Piano non troppo, Plasmic Physics, Pol098, Polonium, Poolkris, Pras, QuackGuru, Quintote, Qxz, RTC, RandomP, Rdsmith4, Remember, Res2216firestar, Retiono Virginian, Rettetast, Reywas92, Reza kalani, Rich Farmbrough, Rifleman 82, Roberta F., RogueNinja, Romanm, Rror, SD6-Agent, SEWilco, SPUI, Samuel, Sander123, Saperaud, Schneelocke, Sengkang, Sephiroth BCR, Seven of Nine, SeventyThree, Sgreddin, Shimmin, Shoefly, Shoeofdeath, Shootbamboo, Skatebiker, SkyLined, Sl, Slowking Man, Smaliali, Smalljim, Smokefoot, Snigbrook, Soroush83, Squids and Chips, Stefkels, Stifynsemons, Stone, Stux, Suisui, Sunborn, Sunbunny, Switchercat, Symplectic Map, Syp, THEN WHO WAS PHONE?, TUF-KAT, Tad Lincoln, Tar7arus, Tealwisp, Tech00, Tetracube, TheBendster, Thricecube, TiCPU, Tim Starling, Tomj, Topbanana, Trumpet marietta 45750, Turian, Typhonius, Until It Sleeps, Urhixidur, VASANTH S.N., VegKilla, Verode, Vivio Testarossa, Vsmith, Warut, Watch37264, Wimt, Winney123, Wknight94, X!, Yath, Yekrats, 590 anonymous edits
Image Sources, Licenses and Contributors image:Br-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Br-TableImage.png License: GNU Free Documentation License Contributors: user:Schneelocke Image: Br,35.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Br,35.jpg License: GNU Free Documentation License Contributors: User:RTC File:Bromine_vial_in_acrylic_cube.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Bromine_vial_in_acrylic_cube.jpg License: unknown Contributors: User:Alchemist-hp Image:Bromine.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Bromine.jpg License: Public Domain Contributors: User:Greenhorn1 Image:Bromine - world production trend.svg Source: http://en.wikipedia.org/w/index.php?title=File:Bromine_-_world_production_trend.svg License: Public Domain Contributors: User:Leyo Image:STS028-96-65.jpg Source: http://en.wikipedia.org/w/index.php?title=File:STS028-96-65.jpg License: Public Domain Contributors: NASA Image:N-Bromosuccinimide Structure.png Source: http://en.wikipedia.org/w/index.php?title=File:N-Bromosuccinimide_Structure.png License: GNU Free Documentation License Contributors: Edgar181, Rhadamante, ~K Image:Tetrabromobisphenol A.svg Source: http://en.wikipedia.org/w/index.php?title=File:Tetrabromobisphenol_A.svg License: Public Domain Contributors: User:Leyo Image:Bromomethane-3D-vdW.png Source: http://en.wikipedia.org/w/index.php?title=File:Bromomethane-3D-vdW.png License: Public Domain Contributors: Benjah-bmm27, Edgar181, Rhadamante Image:AgarosegelUV.jpg Source: http://en.wikipedia.org/w/index.php?title=File:AgarosegelUV.jpg License: GNU Free Documentation License Contributors: TransControl Image:Tralomethrin.png Source: http://en.wikipedia.org/w/index.php?title=File:Tralomethrin.png License: Public Domain Contributors: User:Edgar181 Image:Purpur-mit-Ausfaerbung.png Source: http://en.wikipedia.org/w/index.php?title=File:Purpur-mit-Ausfaerbung.png License: GNU Free Documentation License Contributors: http://de.wikipedia.org/w/index.php?title=Diskussion:Purpur_%28Farbstoff%29
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
12
Krypton
1
Krypton 36
bromine ← krypton → rubidium
Ar ↑
Kr ↓
Xe Periodic Table - Extended Periodic Table
General Name, Symbol, Number
krypton, Kr, 36
Element category
noble gases
Group, Period, Block
18, 4, p
Appearance
colorless gas
Standard atomic weight
83.798(2) g·mol
Electron configuration
[Ar] 3d
−1
10
Electrons per shell
2
6
4s 4p
2, 8, 18, 8 Physical properties
Phase
gas
Density
(0 °C, 101.325 kPa) 3.749 g/L
Melting point
115.79 K (-157.36 °C, -251.25 °F)
Boiling point
119.93 K (-153.22 °C, -244.12 °F)
Triple point
115.775 K, 73.2 kPa
Critical point
209.41 K, 5.50 MPa
Heat of fusion
1.64 kJ·mol−1
Heat of vaporization
9.08 kJ·mol−1
Specific heat capacity
(25 °C) 20.786 J·mol−1·K−1
[1]
Vapor pressure P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
59
65
74
84
99
120
Atomic properties Crystal structure
cubic face centered
Oxidation states
4,
[2]
2
Krypton
2
Electronegativity
3.00 (Pauling scale)
Ionization energies (more)
1st: 1350.8 kJ·mol−1 2nd: 2350.4 kJ·mol−1 3rd: 3565 kJ·mol−1
Covalent radius
116±4 pm
Van der Waals radius
202 pm Miscellaneous [3]
Magnetic ordering
diamagnetic
Thermal conductivity
(300 K) 9.43x10-3 W·m−1·K−1
Speed of sound
(gas, 23 °C) 220 m/s
Speed of sound
(liquid) 1120 m/s
CAS registry number
7439-90-9 Most-stable isotopes
Main article: Isotopes of krypton iso 78
Kr
79
Kr
NA 0.35% syn
half-life 20
2.3×10
y
35.04 h
DM
Kr
81
Kr
2.25% syn
-
78
ε
-
79
β
0.604
79
γ
0.26, 0.39, 0.60
5
2.29×10 y
82
83
11.5%
83
84
57%
84
Kr
85
Kr
86
Kr
syn 17.3%
Br Br
-
Kr is stable with 44 neutron
11.6%
Kr
Se
80
82
Kr
DP
εε
+
80
DE (MeV)
ε
-
γ
0.281
-
0.687
85
81
Br
Kr is stable with 46 neutron Kr is stable with 47 neutron Kr is stable with 48 neutron
10.756 y
β−
Rb
86
Kr is stable with 50 neutron References
Krypton (pronounced /ˈkrɪptɒn/; from Greek: kryptos "hidden") is a chemical element with the symbol Kr and atomic number 36. It is a member of Group 18 and Period 4 elements. A colorless, odorless, tasteless noble gas, krypton occurs in trace amounts in the atmosphere, is isolated by fractionally distilling liquified air, and is often used with other rare gases in fluorescent lamps. Krypton is inert for most practical purposes. Krypton can also form clathrates with water when atoms of it are trapped in a lattice of the water molecules.
Krypton
3
Krypton, like the other noble gases, can be used in lighting and photography. Krypton light has a large number of spectral lines, and krypton's high light output in plasmas allows it to play an important role in many high-powered gas lasers, which pick out one of the many spectral lines to amplify. There is also a specific krypton fluoride laser. The high power and relative ease of operation of krypton discharge tubes caused (from 1960 to 1983) the official meter to be defined in terms of one orange-red spectral line of krypton-86.
Physical properties Krypton is characterized by a brilliant green and orange spectral signature. It is one of the products of uranium fission.[4] Solidified krypton is white and crystalline with a face-centered cubic crystal structure, which is a common property of all noble gases.
History Krypton was discovered in Britain in 1898 by Sir William Ramsay a Scottish chemist and Morris Travers an English chemist in residue left from evaporating nearly all components of liquid air. Neon was discovered by a similar procedure by the same workers just a few weeks later.[5] William Ramsay was awarded the 1904 Nobel Prize in Chemistry for discovery of a series of noble gases, including krypton. A krypton-filled discharge tube in the shape of the element's atomic symbol.
Metric role In 1960, an international agreement defined the meter in terms of wavelength of light emitted by the krypton-86 isotope. This agreement replaced the longstanding standard meter located in Paris, which was a metal bar made of a platinum-iridium alloy (the bar was originally estimated to be one ten-millionth of a quadrant of the earth's polar circumference), and was itself replaced by a definition based on the speed of light — a fundamental physical constant. In October 1983, the Bureau International des Poids et Mesures (International Bureau of Weights and Measures) defined the meter as the distance that light travels in a vacuum during 1/299,792,458 s.[6]
Occurrence The Earth has retained all of the noble gases that were present at its formation except for helium. Helium atoms are very light, and move fast enough to escape the Earth's gravity readily.[7] Krypton's concentration in the atmosphere is about 1 ppm. It can be extracted from liquid air by fractional distillation.[8] The amount of krypton in space is uncertain, as is the amount is derived from the meteoric activity and that from solar winds. The first measurements suggest an overabundance of krypton in space.[9]
Krypton
Compounds Like the other noble gases, krypton is chemically unreactive. However, following the first successful synthesis of xenon compounds in 1962, synthesis of krypton difluoride (KrF2) was reported in 1963.[10] There are unverified reports of other fluorides and a salt of a krypton oxoacid. ArKr+ and KrH+ molecule-ions have been investigated and there is evidence for KrXe or KrXe+.[11] Compounds with krypton bonded to atoms other than fluorine have also been discovered. The reaction of KrF2 with B(OTeF5)3 produces an unstable compound, Kr(OTeF5)2, that contains a krypton-oxygen bond. A krypton-nitrogen bond is found in the cation [HC≡N–Kr–F]+, produced by the reaction of KrF2 with [HC≡NH]+[AsF−6] below −50 °C.[12] At the University of Helsinki in Finland, HKrCN and HKrCCH (krypton hydride-cyanide and hydrokryptoacetylene) were synthesized and determined to be stable up to 40 K.[10]
Isotopes There are 20 known isotopes of krypton.[13] Naturally occurring krypton is made of five stable and one slightly radioactive isotope. Its spectral signature can be produced with some very sharp lines. 81Kr, the product of atmospheric reactions is produced with the other naturally occurring isotopes of krypton. Being radioactive it has a half-life of 230,000 years. Krypton is highly volatile when it is near surface waters but 81Kr has been used for dating old (50,000 - 800,000 year) groundwater.[14] 85
Kr is an inert radioactive noble gas with a half-life of 10.76 years. It is produced by the fission of uranium and plutonium, such as in nuclear bomb testing and nuclear reactors. 85 Kr is released during the reprocessing of fuel rods from nuclear reactors. Concentrations at the North Pole are 30% higher than at the South Pole as most nuclear reactors are in the northern hemisphere.[15]
Applications Krypton's multiple emission lines make ionized krypton gas discharges appear whitish, which in turn makes krypton-based bulbs useful in photography as a brilliant white light source. Krypton is thus used in some types of photographic flashes used in high speed photography. Krypton gas is also combined with other gases to make luminous signs that glow with a bright greenish-yellow light.[16] Krypton mixes with argon as the fill gas of energy saving fluorescent lamps. This reduces their power consumption. Unfortunately this also reduces their light output and raises their cost.[17] Krypton costs 100 times as much as argon. Krypton (along with xenon) is also used to fill incandescent lamps to reduce filament evaporation and allow higher operating temperatures to be used for the filament.[18] A brighter light results which contains more blue than conventional lamps. Krypton's white discharge is often used to good effect in colored gas discharge tubes, which are then simply painted or stained in other ways to allow the desired colour (for example, "neon" type advertising signs where the letters appear in differing colours, are often entirely krypton-based). Krypton is also capable of much higher light power density than neon in the red spectral line region, and for this reason, red lasers for high-power laser light-shows are often krypton lasers with mirrors which select out the red spectral line for laser amplification and emission, rather than the more familiar helium-neon variety, which
4
Krypton
5
could never practically achieve the multi-watt red laser light outputs needed for this application.[19] Krypton has an important role in production and usage of the krypton fluoride laser. The laser has been important in the nuclear fusion energy research community in confinement experiments. The laser has high beam uniformity, short wavelength, and the ability to modify the spot size to track an imploding pellet.[20] In experimental particle physics, liquid krypton is used to construct quasi-homogeneous electromagnetic calorimeters. A notable example is the calorimeter of the NA48 experiment at CERN containing about 27 tons of liquid krypton. This usage is rare, since the cheaper liquid argon is typically used. The advantage of krypton over argon is a small Molière radius of 4.7cm, which allows for excellent spatial resolution and low degree of overlapping. The other parameters relevant for calorimetry application are: radiation length of cm, density of 2.4g/cm³. The sealed spark gap assemblies contained in ignition exciters used in some older jet engines contain a very small amount of Krypton-85 to obtain consistent ionization levels and uniform operation. Krypton-83 has application in magnetic resonance imaging (MRI) for imaging airways. In particular, it may be used to distinguish between hydrophobic and hydrophilic surfaces containing an airway.[21] Although xenon has potential for use in computed tomography (CT) to assess regional ventilation, its anesthetic properties limit its fraction in the breathing gas to 35%. The use of a breathing mixture containing 30% xenon and 30% krypton is comparable in effectiveness for CT to a 40% xenon fraction, while avoiding the unwanted effects of a high fraction xenon gas.[22]
Precautions Very little is known about whether krypton is dangerous. It is considered to be a non-toxic asphyxiant.[23] Krypton has a narcotic potency seven times greater than air, so breathing a gas containing 50% krypton and 50% air would cause narcosis similar to breathing air at four times atmospheric pressure. This would be comparable to scuba diving at a depth of 30 m (100 ft) (see nitrogen narcosis) and potentially could affect anyone breathing it. Nevertheless, that mixture would contain only 10% oxygen and hypoxia would be a greater concern.
Further reading • Los Alamos National Laboratory - Krypton [24] • "Chemical Elements: From Carbon to Krypton" By: David Newton & Lawrence W. Baker • "Krypton 85: a Review of the Literature and an Analysis of Radiation Hazards" By: William P. Kirk.
External links • WebElements.com – Krypton
[25]
• Krypton Fluoride Lasers [26] • Computational Chemistry Wiki
[27]
Krypton
References [1] "Section 4, Properties of the Elements and Inorganic Compounds; Melting, boiling, triple, and critical temperatures of the elements". CRC Handbook of Chemistry and Physics (85th edition ed.). Boca Raton, Florida: CRC Press. 2005. [2] " Krypton: krypton(IV) fluoride compound data (http:/ / books. google. com/ books?id=ZCkKH4bUcPUC& pg=RA1-PA341& lpg=RA1-PA341& dq="krypton+ iv"& source=web& ots=XA_ev9xdZZ& sig=vla8Z9t5gPpCRw2rm-cXK1Tsgt8)". Books.Google.com. . Retrieved 2007-12-10. [3] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [4] " Krypton (http:/ / www. ead. anl. gov/ pub/ doc/ krypton. pdf)" (in English). Argonne National Laboratory, EVS. 08 2005. pp. 1. . Retrieved 2007-03-17. [5] William Ramsay, Morris W. Travers (1898). " On a New Constituent of Atmospheric Air (http:/ / links. jstor. org/ sici?sici=0370-1662(1898)63<405:OANCOA>2. 0. CO;2-M)". Proceedings of the Royal Society of London 63: 405–408. doi: 10.1098/rspl.1898.0051 (http:/ / dx. doi. org/ 10. 1098/ rspl. 1898. 0051). . [6] Gibbs, Philip (1997). " How is the speed of light measured? (http:/ / math. ucr. edu/ home/ baez/ physics/ Relativity/ SpeedOfLight/ measure_c. html)" (in English). Department of Mathematics, University of California. . Retrieved 2007-03-19. [7] Escape of Gases from the Atmosphere [8] " How Products are Made: Krypton (http:/ / www. madehow. com/ Volume-4/ Krypton. html)". . Retrieved 2006-07-02. [9] Cardelli, Jason A.; Meyer, David M. (18). " The Abundance of Interstellar Krypton (http:/ / www. journals. uchicago. edu/ doi/ full/ 10. 1086/ 310513)" (in English). The Astrophysical Journal Letters. The American Astronomical Society. pp. L57–L60. . Retrieved 2007-04-05. [10] Bartlett, Neil (2003). " The Noble Gases (http:/ / pubs. acs. org/ cen/ 80th/ noblegases. html)" (in English). Chemical & Engineering News. . Retrieved 2006-07-02. [11] " Periodic Table of the Elements (http:/ / www. bu. edu/ ehs/ ih/ pdf/ periodic_table. pdf)" (in English). Los Alamos National Laboratory's Chemistry Division. pp. 100-101. . Retrieved 2007-04-05. [12] John H. Holloway; Eric G. Hope (1998). A. G. Sykes. ed. Advances in Inorganic Chemistry. Academic Press. pp. 57. ISBN 012023646X. [13] " Isotopes of Krypton (http:/ / ie. lbl. gov/ education/ parent/ Kr_iso. htm)". Nuclear Science Division. . Retrieved 2007-03-20. [14] Thonnard, Norbert; Larry D. MeKay, Theodore C. Labotka (31). " Development of Laser-Based Resonance Ionization Techniques for 81-Kr and 85-Kr Measurements in the Geosciences (http:/ / www. osti. gov/ bridge/ servlets/ purl/ 809813-0zMCV1/ native/ 809813. pdf)" (in English). University of Tennessee, Institute for Rare Isotope Measurements. pp. 4-7. . Retrieved 2007-03-20. [15] " Resources on Isotopes (http:/ / wwwrcamnl. wr. usgs. gov/ isoig/ period/ kr_iig. html)". U.S. Geological Survey. . Retrieved 2007-03-20. [16] " Mercury in Lighting (http:/ / www. capecodextension. org/ pdfs/ Mercury Lighting. pdf)". Cape Cod Cooperative Extension. . Retrieved 2007-03-20. [17] "Energy-saving" lamps (http:/ / www. anaheim. net/ utilities/ ea/ PA_11. html) [18] Properties, Applications and Uses of the "Rare Gases" Neon, Krypton and Xenon (http:/ / www. uigi. com/ rare_gases. html) [19] " Laser Devices, Laser Shows and Effect (http:/ / www. gameops. com/ content/ pdf/ laser_terms. pdf)" (PDF). . Retrieved 2007-04-05. [20] Sethian, J.; M. Friedman, M.Myers. " Krypton Fluoride Laser Development for Inertial Fusion Energy (http:/ / aries. ucsd. edu/ LIB/ MEETINGS/ IAEAIFECRP/ PDF/ Sethian. pdf)" (in English). Plasma Physics Division, Naval Research Laboratory. pp. 1-8. . Retrieved 2007-03-20. [21] Pavlovskaya, GE; Cleveland, ZI; Stupic, KF; Basaraba, RJ; Meersmann, T (December 2005). " Hyperpolarized krypton-83 as a contrast agent for magnetic resonance imaging (http:/ / www. pnas. org/ content/ 102/ 51/ 18275. abstract)". Proceedings of the National Academy of Sciences U.S.A. 102 (51): 18275–9. doi: 10.1073/pnas.0509419102 (http:/ / dx. doi. org/ 10. 1073/ pnas. 0509419102). PMID 16344474. PMC: 1317982 (http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?tool=pmcentrez& artid=1317982). . [22] Chon, D; Beck, KC; Simon, BA; Shikata, H; Saba, OI; Hoffman, EA (April 2007). " Effect of low-xenon and krypton supplementation on signal/noise of regional CT-based ventilation measurements (http:/ / jap. physiology. org/ cgi/ content/ abstract/ 102/ 4/ 1535)". Journal of Applied Physiology 102 (4): 1535–44. doi: 10.1152/japplphysiol.01235.2005 (http:/ / dx. doi. org/ 10. 1152/ japplphysiol. 01235. 2005). PMID 17122371. . [23] http:/ / pt. chemicalstore. com/ Kr%20-%20Krypton. html [24] http:/ / periodic. lanl. gov/ elements/ 36. html
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Krypton [25] http:/ / www. webelements. com/ webelements/ elements/ text/ Kr/ index. html [26] http:/ / other. nrl. navy. mil/ LaserFusionEnergy/ lasercreation. htm [27] http:/ / www. compchemwiki. org/ index. php?title=Krypton
7
Article Sources and Contributors
Article Sources and Contributors Krypton Source: http://en.wikipedia.org/w/index.php?oldid=307571150 Contributors: 2D, 8v, 99of9, A new name 2008, ABF, Aarchiba, Abrech, AbsolutDan, Achaemenes, AeroChemistry, Ahoerstemeier, AirdishStraus, Aka042, Alan012, Alansohn, Ale jrb, AmiDaniel, Amicon, Anakinjmt, Anclation, AndonicO, Andre Engels, AndreasJS, Andrei901, Andres, AngelOfSadness, Animum, Anoop.m, Antandrus, Anthony Appleyard, Antonio Lopez, ApsbaMd2, Archimerged, Astroview120mm, AuburnPilot, Auror, BRG, Badocter, Beetstra, BenBaker, Benbest, Bencherlite, Benjiboi, Bethling, Bigvis, Blanchardb, Blathnaid, BlueEarth, Blueknightex, Bobo192, Bongwarrior, Boyofus, Brideshead, BrucePodger, Bryan Derksen, Bushido Hacks, Buttman12345, C66, CYD, Cameron Nedland, Can't sleep, clown will eat me, CanadianLinuxUser, Captain Wikify, Carnildo, Casper2k3, CharlotteWebb, Chem-awb, Chitrapa, Clark89, Cobi, Con soccer13, Conversion script, Crazy British Kid, Cryptic C62, Crzycheetah, Cst17, Ctjj.stevenson, Cunaaay, Cyrus Andiron, DVD R W, Da monster under your bed, Danno uk, Danski14, Darrien, Darth Panda, David Latapie, Davidhorman, Dcandeto, Deli nk, Delldot, Deltabeignet, Deor, Derek.cashman, Deryck Chan, Deville, Dgrant, Discospinster, DisorderedOne, Dlohcierekim's sock, Donarreiskoffer, Doulos Christos, Dreadstar, Druiddude2384, Dtkey, Dureo, Dysepsion, EJF, ERK, Edgar181, Edsanville, Edward321, Egomaniac, El, El C, Elkman, Ellimist, Emperorbma, Eng02019, Epbr123, Erik Zachte, Excirial, Exert, Fedallah, Femto, Finngall, Fish,chips+mushypeas, Fonzy, Frankenpuppy, Frosty0814snowman, GDonato, GPHemsley, Giftlite, Giggy, GlassCobra, Gogo Dodo, Goudzovski, GraemeL, GreenGourd, Grendelkhan, Guitarlord123, Gurch, Hak-kâ-ngìn, Harland1, Harryboyles, Howieju, Husond, Iago4096, Iapetus, Icairns, Ignatzmice, Imaninjapirate, Imsasn, Indon, Indosauros, Instinct, Iridescent, Itub, J.delanoy, JFreeman, Jaknouse, Jaraalbe, Jason Leach, Jay32183, Jeffrey Mall, Jeffwood12, Jeronimo, Jj137, Jjjsixsix, Jmrowland, Joanjoc, JoanneB, John, John254, Johnbrownsbody, Jojhutton, Jose77, Joshschr, Julesd, Juliancolton, Junglecat, Jvbishop, Kangxi emperor6868, Katalaveno, Kbh3rd, Keegan, Keenman76, Keilana, Kesac, Kevin Breitenstein, Khukri, Killawolf, King of Hearts, Kingpin13, KnowledgeOfSelf, Kowey, Kozuch, Kubigula, Kurykh, Kvdveer, Kwamikagami, LarryMorseDCOhio, LeaveSleaves, Levil, Lexor, Longhair, Lord Emsworth, LordFoppington, LuigiManiac, Lukemynds, Ma335019, Macuser678, Mah213, Malcolm Farmer, Mandarax, MarSch, Marauder40, Marek69, Martin-vogel, MasterXC, Materialscientist, Mattbaileyisneat, MattieTK, Matts computer, Mav, Maximilli, Mecee, Mgimpel, Michael Rawdon, MichaelMaggs, Mike Christie, Mike Winters, Mike s, Minesweeper, Minnesota1, Mixwell, Mr. Wheely Guy, Muke, Mysid, NHRHS2010, Nadando, NawlinWiki, Nergaal, Neurolysis, Ngchen, Nick, Nihiltres, NodnarbLlad, Oliver202, Opabinia regalis, Ossmann, Oxymoron83, Partnerintime12, Pekaje, Pendragonneo, PeterJeremy, Pgk, Pharaoh of the Wizards, Philip Trueman, Physchim62, Piano non troppo, Pifreak94, Plasmic Physics, Polonium, Poolkris, Pplz4ever, Pre1mjr, Proofreader77, Pseudomonas, Pundit, Q43, Quiddity, Quintote, Qwertyasfgh, Qxz, R, RTC, RazorICE, Red Director, Red Thunder, Remember, Reuben, RexNL, RexxS, RichardB, RiddG07, Rixshk, Rlevse, Roberta F., Rodney Ruff, Romanm, Rossdaboss50, Ryanh1994, SJP, SMC, Sannse, Saperaud, Sbharris, Schneelocke, Scientizzle, Scohoust, Sengkang, Shafei, Shakowski, Sheitan, SimonP, Singergirl350, SiriusGrey, Skagedal, Skizzik, Sl, Smalljim, Smallzman, Snowolf, Someguy1221, Somno, Split Infinity, Squids and Chips, Stan J Klimas, Stephen, Stifynsemons, Stismail, Stone, Suisui, Sunborn, Superman1560, THEN WHO WAS PHONE?, TJDay, Tarret, Tcncv, Tellyaddict, Tempodivalse, Tetracube, The Rambling Man, The_ansible, Therealsquee, Thricecube, Thue, TicketMan, Tiddly Tom, Tim Starling, TimBuck2, TimVickers, Tiptoety, Tjmage1, Triforce of Power, Trojancowboy, UberScienceNerd, Uncle Dick, Utcursch, Van helsing, Veemonkamiya, Versus22, Vsmith, Vuong Ngan Ha, WadeSimMiser, Waggers, Warut, Watch37264, Weregerbil, Wideeyedraven, WikiZorro, Wilder22, Will dogson, William Avery, Willking1979, Wknight94, Wrenchelle, Writtenright, X!, XxLayotaxX, Yamamoto Ichiro, Yekrats, Yelyos, Yoshie310, Yyy, ZenerV, Zotel, Александър, 981 anonymous edits
Image Sources, Licenses and Contributors image:Kr-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Kr-TableImage.png License: GNU Free Documentation License Contributors: user:Schneelocke Image:KrTube.jpg Source: http://en.wikipedia.org/w/index.php?title=File:KrTube.jpg License: unknown Contributors: User:Pslawinski
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
8
Rubidium
1
Rubidium 37
krypton ← rubidium → strontium
K ↑
Rb ↓
Cs Periodic Table - Extended Periodic Table
General Name, Symbol, Number
rubidium, Rb, 37
Element category
alkali metals
Group, Period, Block
1, 5, s
Appearance
grey white
Standard atomic weight
85.4678(3) g·mol
Electron configuration
[Kr] 5s
−1
1
Electrons per shell
2, 8, 18, 8, 1 Physical properties
Phase
solid
Density (near r.t.)
1.532 g·cm−3
Liquid density at m.p.
1.46 g·cm−3
Melting point
312.46 K (39.31 °C, 102.76 °F)
Boiling point
961 K (688 °C, 1270 °F)
Critical point
(extrapolated) 2093 K, 16 MPa
Heat of fusion
2.19 kJ·mol−1
Heat of vaporization
75.77 kJ·mol−1
Specific heat capacity
(25 °C) 31.060 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
434
486
552
641
769
958
Rubidium
2 Atomic properties
Crystal structure
body centered cubic
Oxidation states
1 (strongly basic oxide)
Electronegativity
0.82 (Pauling scale)
Ionization energies (more)
1st: 403.0 kJ·mol−1 2nd: 2633 kJ·mol−1 3rd: 3860 kJ·mol−1
Atomic radius
248 pm
Covalent radius
220±9 pm
Van der Waals radius
244 pm Miscellaneous [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(20 °C) 128 n Ω·m
Thermal conductivity
(300 K) 58.2 W·m
Speed of sound (thin rod)
(20 °C) 1300 m/s
Young's modulus
2.4 GPa
Bulk modulus
2.5 GPa
−1
Mohs hardness
−1
·K
0.3
Brinell hardness
0.216 MPa
CAS registry number
7440-17-7 Most-stable isotopes
Main article: Isotopes of rubidium iso 83
Rb
NA syn
half-life 86.2 d
DM ε γ
84
Rb
syn
32.9 d
ε β+
85
Rb
86
Rb
72.168% syn
Rb
27.835%
0.52, 0.53, 0.55
DP 83
Kr
-
-
84
1.66, 0.78
84
Kr Kr
γ
0.881
-
β−
0.892
84
1.775
86
Sr
85
Rb is stable with 48 neutron
18.65 d
β− γ
87
DE (MeV)
4.88 × 1010 y
β− References
1.0767 0.283
Sr
87
Sr
Rubidium
Rubidium (pronounced /rʊˈbɪdiəm/) is a chemical element with the symbol Rb and atomic number 37. Rb is a soft, silvery-white metallic element of the alkali metal group. Rubidium is very soft and highly reactive, with properties similar to other elements in group 1, such as very rapid oxidation in air. Its compounds have chemical and electronic applications. Rubidium metal is easily vaporized and has a convenient spectral absorption range, making it a frequent target for laser manipulation of atoms. Rubidium is not known to be necessary for any living organisms. However, like caesium, rubidium ions are handled by living organisms in a manner similar to potassium: it is actively taken up by plants and by living animals' cells. Rubidium has one stable isotope,85Rb. The isotope 87Rb which composes almost 28% of naturally occurring rubidium is slightly radioactive, with a half-life of 49 billion years—more than three times longer than the estimated age of the universe.
Characteristics Rubidium is the second most electropositive of the stable alkali elements and liquefies at a high ambient temperature, 39.3 °C (102.7 °F). Like other group 1 elements this metal reacts violently in water. In common with potassium and caesium this reaction is usually vigorous enough to ignite the liberated hydrogen. Rubidium has also been reported to ignite spontaneously in air. Also like other alkali metals, it forms amalgams with mercury and it can form alloys with gold, caesium, sodium, and potassium. The element gives a reddish-violet color to a flame, hence its name.
History Rubidium (L rubidus, deepest red) was discovered in 1861 by Robert Bunsen and Gustav Kirchhoff in the mineral lepidolite through the use of a spectroscope.[2] The extraction of 150 kg of lepidolite yielded only a few grams for analysis. The first rubidium metal was produced by the reaction of rubidium chloride with potassium by Bunsen.
Occurrence Rubidium is about the twenty-third[3] most abundant element in the Earth's crust, roughly as abundant as zinc and rather more common than copper. It occurs naturally in the minerals leucite, pollucite, carnallite and zinnwaldite, which contains traces of up to 1% of its oxide. Lepidolite contains 1.5% rubidium and this is the commercial source of the element. Some potassium minerals and potassium chlorides also contain the element in commercially significant amounts. One notable source is also in the extensive deposits of pollucite at Bernic Lake, Manitoba (also a source of the related element caesium). Rubidium metal can be produced by reducing rubidium chloride with calcium among other methods. In 1997 the cost of this metal in small quantities was about US$25/gram.
3
Rubidium
Isotopes There are 26 isotopes of rubidium known with naturally occurring rubidium being composed of just two isotopes; Rb-85 (72.2%) and the radioactive Rb-87 (27.8%). Natural rubidium is radioactive with specific activity of about 670 Bq/g, enough to fog photographic film in approximately 30 to 60 days. Rb-87 has a half-life of 4.88 × 1010 years. It readily substitutes for potassium in minerals, and is therefore fairly widespread. Rb has been used extensively in dating rocks; Rb-87 decays to stable strontium-87 by emission of a negative beta particle. During fractional crystallization, Sr tends to become concentrated in plagioclase, leaving Rb in the liquid phase. Hence, the Rb/Sr ratio in residual magma may increase over time, resulting in rocks with increasing Rb/Sr ratios with increasing differentiation. Highest ratios (10 or higher) occur in pegmatites. If the initial amount of Sr is known or can be extrapolated, the age can be determined by measurement of the Rb and Sr concentrations and the Sr-87/Sr-86 ratio. The dates indicate the true age of the minerals only if the rocks have not been subsequently altered. See Rubidium-Strontium dating for a more detailed discussion.
Uses and applications Rubidium had minimal industrial use until the 1930s. Historically, the most important use for rubidium has been in research and development, primarily in chemical and electronic applications. In 1999 rubidium-87 was used to make a Bose-Einstein condensate[4] , for which the discoverers won the 2001 Nobel Prize in Physics[5] . Rubidium is easily ionized, so it has been considered for use in ion engines for space vehicles (but caesium and xenon are more efficient for this purpose). Rubidium compounds are sometimes used in fireworks to give them a purple color. RbAg4I5 has the highest room temperature conductivity of any known ionic crystal. This property could be useful in thin film batteries and in other applications.[6] Rubidium has also been considered for use in a thermoelectric generator using the magnetohydrodynamic principle, where rubidium ions are formed by heat at high temperature and passed through a magnetic field. These conduct electricity and act like an armature of a generator thereby generating an electric current. Rubidium, particularly 87Rb, in the form of vapor, is one of the most commonly used atomic species employed for laser cooling and Bose-Einstein condensation. Its desirable features for this application include the ready availability of inexpensive diode laser light at the relevant wavelength, and the moderate temperatures required to obtain substantial vapor pressures. Rubidium has been used for polarizing 3He (that is, producing volumes of magnetized 3He gas, with the nuclear spins aligned toward a particular direction in space, rather than randomly). Rubidium vapor is optically pumped by a laser and the polarized Rb polarizes 3 He by the hyperfine interaction.[7] Spin-polarized 3He cells are becoming popular for neutron polarization measurements and for producing polarized neutron beams for other purposes.[8] Rubidium is the primary compound used in secondary frequency references (Rubidium Oscillators) to maintain frequency accuracy in cell site transmitters and other electronic
4
Rubidium transmitting, networking and test equipment. Rubidium references are often used with GPS to produce a "Primary Frequency Standard" that has greater accuracy but is less expensive than Cesium standards. Rubidium references such as the LPRO series from Datum were mass-produced for the Telecom industry and are now selling for under one hundred dollars on the secondary market. The general life expectancy is 10 years or better for most designs. Other potential or current uses of rubidium include: • A working fluid in vapor turbines. • A getter in vacuum tubes. • A photocell component. • The resonant element in atomic clocks. This is due to the hyperfine structure of rubidium's energy levels. • An ingredient in special types of glass. • The production of superoxide by burning in oxygen. • The study of potassium ion channels in biology. • Rubidium is used to locate brain tumours, due to its slight radioactivity. [9] • Rubidium vapor has been used to make atomic magnetometers. 87Rb is currently being used, with other alkali metals, in the development of spin-exchange relaxation-free (SERF) magnetometers.[10]
Compounds Rubidium chloride is probably the most-used rubidium compound; it is used in biochemistry to induce cells to take up DNA, and as a biomarker since it is readily taken up to replace potassium, and does not normally occur in living organisms. Rubidium hydroxide is the starting material for most rubidium-based chemical processes; rubidium carbonate is used in some optical glasses. Rubidium has a number of oxides, including Rb6O and Rb9O2 which appear if rubidium metal is left exposed to air; the final product of reacting with oxygen is the superoxide RbO2. Rubidium forms salts with most anions. Some common rubidium compounds are rubidium chloride (RbCl), rubidium monoxide (Rb2O) and rubidium copper sulfate Rb2SO4·CuSO4·6H2O). A compound of rubidium, silver and iodine, RbAg4I5, has interesting electrical characteristics and might be useful in thin film batteries.[11]
Precautions Rubidium reacts violently with water and can cause fires. To ensure both health and safety and purity, this element must be kept under a dry mineral oil, and in practice is usually sealed in glass ampules in an inert atmosphere. Rubidium forms peroxides on exposure to even air diffusing into oil, and is thus subject to some of the same peroxide precautions as storage of metallic potassium.
5
Rubidium
6
Biological effects Rubidium, like sodium and potassium, is almost always in its +1 oxidation state when dissolved in water, and this includes all biological systems. The human body tends to treat Rb+ ions as if they were potassium ions, and therefore concentrates rubidium in the body's intracellular fluid (i.e., inside cells). The ions are not particularly toxic, and are relatively quickly removed in the sweat and urine. As a result of changes in the blood brain barrier in brain tumors, rubidium collects more in brain tumors than normal brain tissue, allowing short-lived radioisotopes of rubidium to be used in nuclear medicine to locate and image brain tumors.
Sources • Los Alamos National Laboratory – Rubidium
[12]
• Louis Meites, Handbook of Analytical Chemistry (New York: McGraw-Hill Book Company, 1963) • Daniel A. Steck. "Rubidium-87 D Line Data [13]". Los Alamos National Laboratory (technical report LA-UR-03-8638). http:/ / george. ph. utexas. edu/ ~dsteck/ alkalidata/ rubidium87numbers. pdf.
External links • WebElements.com – Rubidium
[14]
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] G. Kirchhoff, R. Bunsen (1861). "Chemische Analyse durch Spectralbeobachtungen". Annalen der Physik und Chemie 189 (7): 337–381. doi: 10.1002/andp.18611890702 (http:/ / dx. doi. org/ 10. 1002/ andp. 18611890702). [3] http:/ / www. rsc. org/ chemsoc/ visualelements/ pages/ data/ rubidium_data. html [4] " Bose-Einstein Condensation (http:/ / www. bookrags. com/ BoseâEinstein_condensate)". World of Physics on Bose-Einstein Condensation. BookRags. . Retrieved 2008-01-26. [5] Levi, Barbara Goss (2001). " Cornell, Ketterle, and Wieman Share Nobel Prize for Bose-Einstein Condensates (http:/ / www. physicstoday. org/ pt/ vol-54/ iss-12/ p14. html)". Search & Discovery. Physics Today online. . Retrieved 2008-01-26. [6] Bradley, J. N.; Greene, P. D. (1967). "Relationship of structure and ionic mobility in solid MAg4I5". Trans. Faraday Soc. 63: 2516. doi: 10.1039/TF9676302516 (http:/ / dx. doi. org/ 10. 1039/ TF9676302516). [7] Gentile, T.R. et al.. " Polarized 3He spin filters for slow neutron physics (http:/ / nvl. nist. gov/ pub/ nistpubs/ jres/ 110/ 3/ j110-3gen. pdf)". Journal of Research of the National Institute of Standards and Technology 100: 299. . [8] " Neutron spin filters based on polarized helium-3 (http:/ / www. ncnr. nist. gov/ AnnualReport/ FY2002_html/ pages/ neutron_spin. htm)". NIST Center for Neutron Research 2002 Annual Report. . [9] http:/ / www. rsc. org/ chemsoc/ visualelements/ pages/ data/ rubidium_data. html [10] Li, Zhimin et al. (2006). " Parametric modulation of an atomic magnetometer (http:/ / scitation. aip. org/ getabs/ servlet/ GetabsServlet?prog=normal& id=APPLAB000089000013134105000001& idtype=cvips& gifs=yes)". Applied Physics Letters 89: 134105. doi: 10.1063/1.2357553 (http:/ / dx. doi. org/ 10. 1063/ 1. 2357553). . [11] Smart, Lesley; Moore, Elaine (1995). " RbAg4I5 (http:/ / books. google. de/ books?id=pVw98i6gtwMC& pg=PA176)". Solid state chemistry: an introduction. CRC Press. pp. 176–177. ISBN 9780748740680. . [12] http:/ / periodic. lanl. gov/ elements/ 37. html [13] http:/ / george. ph. utexas. edu/ ~dsteck/ alkalidata/ rubidium87numbers. pdf [14] http:/ / www. webelements. com/ webelements/ elements/ text/ Rb/ index. html
Article Sources and Contributors
Article Sources and Contributors Rubidium Source: http://en.wikipedia.org/w/index.php?oldid=306311243 Contributors: 65.68.87.xxx, A2Kafir, Ahoerstemeier, Alansohn, Ale jrb, Andre Engels, AndreasJS, Andres, Anwar saadat, Archimerged, Ariel Pontes, Arkania, Arkuat, Beatnikbob, Benbest, Bender235, BlueEarth, Bobo192, Bogdangiusca, Bonzostar, Bryan Derksen, CRGreathouse, CYD, Can't sleep, clown will eat me, Capricorn42, Carnildo, Celtic-Boy-007, ChicXulub, Colbuckshot, Conversion script, Cosmium, Dajagr, Darrien, David Latapie, Dead3y3, Deglr6328, Delta G, Derek Ross, Discospinster, Dnn87, Doktor, Donarreiskoffer, DrBob, E rulez, Edgar181, El C, Elendil's Heir, Ellmist, Emperorbma, Epbr123, Eric-Wester, Erik Zachte, Exmophead, Femto, Fivemack, Fizzy, Flewis, Fonzy, Foonly, Gail, Gbr3, Gcsuchemistry, Giftlite, GorillaWarfare, Grendelkhan, Hadal, Hak-kâ-ngìn, Heron, IRKAIN, Icairns, Ideyal, It Is Me Here, Ixfd64, J.delanoy, JForget, Jamesbateman, Jaraalbe, Jeffrey O. Gustafson, Jeronimo, John, John254, Jose77, Karl-Henner, Karlhahn, Katalaveno, Khargas, Kindt, Kingpin13, Kwamikagami, LAX, LeaveSleaves, Lectonar, Levil, Lindmere, Llamadog903, Loict, LuigiManiac, Luk, Mako098765, Malbi, Malcolm Farmer, Martarius, Mav, Mervyn, Mgimpel, Minesweeper, Mr0t1633, Munboy, NFreak007, Natalie Erin, Nergaal, Nihiltres, Nikai, Omegatron, Philip Trueman, Piano non troppo, Plexust, Polonium, Poolkris, Poopy45, Pras, Quadell, Quercus basaseachicensis, RJF19, RJaguar3, RTC, Rambam rashi, Rd232, Red King, Remember, Riana, Romanm, Rsocol, Saperaud, Sbharris, Schneelocke, Sengkang, Shirulashem, SimonP, Sinneed, Sirex98, Skatebiker, Sl, Smalljim, Squids and Chips, Srtxg, Srushe, Stan J Klimas, StaticGull, Stifynsemons, Stone, Sunborn, Svante, Synchronism, Tagishsimon, Tantalate, Tapir Terrific, Tarquin, Tempodivalse, Thingg, Thricecube, Tim Starling, Timeineurope, Toddst1, Tresiden, UkPaolo, Utcursch, UtherSRG, Van helsing, Velvetron, Versus22, Vsmith, Vuo, Vuong Ngan Ha, Whosasking, Whydoyoucarewhatmynameis, Wilbern Cobb, Wimt, Yekrats, YukoValis, Zara1709, 308 anonymous edits
Image Sources, Licenses and Contributors image:Rb-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Rb-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Mav, Paddy, Saperaud Image:Rb66.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Rb66.jpg License: unknown Contributors: Dnn87, Radiant chains, 13 anonymous edits
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
7
Strontium
1
Strontium 38
rubidium ← strontium → yttrium
Ca ↑
Sr ↓
Ba Periodic Table - Extended Periodic Table
General Name, Symbol, Number
strontium, Sr, 38
Element category
alkaline earth metals
Group, Period, Block
2, 5, s
Appearance
silvery white metallic
Standard atomic weight
87.62(1) g·mol
Electron configuration
[Kr] 5s
−1
2
Electrons per shell
2, 8, 18, 8, 2 Physical properties
Phase
solid
Density (near r.t.)
2.64 g·cm−3
Liquid density at m.p.
2.375 g·cm−3
Melting point
1050 K (777 °C, 1431 °F)
Boiling point
1655 K (1382 °C, 2520 °F)
Heat of fusion
7.43 kJ·mol−1
Heat of vaporization
136.9 kJ·mol−1
Specific heat capacity
(25 °C) 26.4 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
796
882
990
1139
1345
1646
Atomic properties
Strontium
2
Crystal structure
cubic face centered
Oxidation states
2, 1 (strongly basic oxide)
Electronegativity
0.95 (Pauling scale)
[1]
Ionization energies (more)
1st: 549.5 kJ·mol−1 2nd: 1064.2 kJ·mol−1 3rd: 4138 kJ·mol−1
Atomic radius
215 pm
Covalent radius
195±10 pm Miscellaneous
Magnetic ordering
paramagnetic
Electrical resistivity
(20 °C) 132 n Ω·m
Thermal conductivity
(300 K) 35.4 W·m
Thermal expansion
(25 °C) 22.5 µm·m
Shear modulus
6.1 GPa
−1
−1
Poisson ratio
0.28
Mohs hardness
1.5
CAS registry number
7440-24-6
−1
·K
−1
·K
Most-stable isotopes
Main article: Isotopes of strontium iso
NA
half-life
82
syn
25.36 d
83
syn
1.35 d
Sr Sr
84
Sr
85
Sr
0.56% syn
DM ε
82
-
83
β+
1.23
83
γ
0.76, 0.36
ε
64.84 d
86
87
7.0%
87
88
82.58%
88
89
Sr
90
Sr
syn
syn
Rb Rb
-
Sr is stable with 46 neutron
9.86%
Sr
Rb
84
86
Sr
DP
-
ε γ
Sr
DE (MeV)
0.514D
85
Rb
-
Sr is stable with 48 neutron Sr is stable with 49 neutron Sr is stable with 50 neutron
50.52 d
28.90 y
ε
1.49
89
β−
0.909D
89
β−
0.546
90
Rb Y Y
Strontium
3 References
Strontium (pronounced /ˈstrɒnʃiəm/, /ˈstrɒntiəm/, or /ˈstrɒnʃəm/) is a chemical element with the symbol Sr and the atomic number 38. An alkaline earth metal, strontium is a soft silver-white or yellowish metallic element that is highly reactive chemically. The metal turns yellow when exposed to air. It occurs naturally in the minerals celestine and strontianite. The 90Sr isotope is present in radioactive fallout and has a half-life of 28.90 years. Both strontium and strontianite are named after Strontian, a village in Scotland near which the mineral was first discovered.
Characteristics Due to its extreme reactivity with oxygen and water, this element occurs naturally only in compounds with other elements, as in the minerals strontianite and celestite. Strontium is a grey/silvery metal that is softer than calcium and even more reactive in water, with which strontium reacts on contact to produce strontium hydroxide and hydrogen gas. It burns in air to produce both strontium oxide and strontium nitride, but since it does not react with nitrogen below 380°C it will only form the oxide spontaneously at room temperature. It should be kept under kerosene to prevent oxidation; freshly exposed strontium metal rapidly turns a yellowish color with the formation of the oxide. Finely powdered strontium metal will ignite spontaneously in Dendritic oxidized strontium air at room temperature. Volatile strontium salts impart a crimson color to flames, and these salts are used in pyrotechnics and in the production of flares. Natural strontium is a mixture of four radiostable isotopes.
Compounds • Ferrite magnets and refining zinc. • Strontium titanate has an extremely high refractive index and an optical dispersion greater than that of diamond, making it useful in a variety of optics applications. This quality has also led to it being cut into gemstones, in particular as a diamond simulant. However, it is very soft and easily scratches so it is rarely used. • Strontium carbonate, strontium nitrate, and strontium sulfate are commonly used in fireworks for red color. • Strontium aluminate is used as a bright phosphor with long persistence of phosphorescence. • Strontium chloride is sometimes used in toothpastes for sensitive teeth. One popular brand includes 10% total strontium chloride hexahydrate by weight. • Strontium oxide is sometimes used to improve the quality of some pottery glazes.
Strontium
4
• Strontium ranelate is used in the treatment of osteoporosis. It is a prescription drug in the EU, but not in the USA. • Strontium barium niobate is used in large scale outdoors Holgraphic displays as a "screen". When projected with blue laser light images are 3-D and very realistic.
Isotopes Strontium has four stable, naturally occurring isotopes: 84Sr (0.56%), 86Sr (9.86%), 87Sr (7.0%) and 88Sr (82.58%). Only 87Sr is radiogenic; it is produced by decay from the radioactive alkali metal 87Rb, which has a half-life of 4.88 × 1010 years. Thus, there are two sources of 87Sr in any material: that formed in stars along with 84Sr, 86Sr and 88Sr, as well as that formed by radioactive decay of 87Rb. The ratio 87Sr/86Sr is the parameter typically reported in geologic investigations; ratios in minerals and rocks have values ranging from about 0.7 to greater than 4.0. Because strontium has an atomic radius similar to that of calcium, it readily substitutes for Ca in minerals. Sixteen unstable isotopes are known to exist. Of greatest importance are half-life of 28.78 years and 89Sr with a half-life of 50.5 days. •
90
90
Sr with a
Sr is a by-product of nuclear fission which is found in nuclear fallout and presents a health problem since it substitutes for calcium in bone, preventing expulsion from the body. This isotope is one of the best long-lived high-energy beta emitters known, and is used in SNAP (Systems for Nuclear Auxiliary Power) devices. These devices hold promise for use in spacecraft, remote weather stations, navigational buoys, etc, where a lightweight, long-lived, nuclear-electric power source is required. The 1986 Chernobyl nuclear accident contaminated a vast area with 90Sr. 90Sr confined inside a concave silver plaque is also used for the medical treatment of a resected pterygium. • 89Sr is a short-lived artificial radioisotope which is used in the treatment of bone cancer. In circumstances where cancer patients have widespread and painful bony metastases (secondaries), the administration of 89Sr results in the delivery of radioactive emissions (beta particles in this case) directly to the area of bony problem (where calcium turnover is greatest). The 89Sr is manufactured as the chloride salt (which is soluble), and when dissolved in normal saline can be injected intravenously. Typically, cancer patients will be treated with a dose of 150 MBq. The patient needs to take precautions following this because their urine becomes contaminated with radioactivity, so they need to sit to urinate and double flush the toilet. The beta particles travel about 3.5mm in bone (energy 0.583 MeV) and 6.5mm in tissue, so there is no requirement to isolate patients who have been treated except to say they should not have any one (especially young children) sitting in their laps for 10-40 days. The variation in time results from the variable clearing time for 89Sr which depends on renal function and the number of bony metastases. With a lot of bony metastases, the entire 89Sr dose can be taken up into bone and so the entire radioactivity is retained to decay over a 50.5 day half-life. However, where there are few bony metastases, the large proportion of 89Sr not taken up by the bone will be filtered by the kidney, so that the effective half-life (a combination of the physical and biological half-life) will be much shorter.
Strontium
5
History The mineral strontianite is named after the Scottish village of Strontian, having been discovered in the lead mines there in 1787.[2] Adair Crawford recognized it as differing from other barium minerals in 1790. Strontium itself was discovered in 1798 by Thomas Charles Hope, and metallic strontium was first isolated by Sir Humphry Davy in 1808 using electrolysis of a mixture containing strontium chloride and mercuric oxide and announced by him in a lecture to the Royal Society on 30 June 1808[3] .
Occurrence According to the British Geological Survey, China was the top producer of strontium in 2007, with over two-thirds world share, followed by Spain and Mexico.[4] Strontium commonly occurs in nature, the 15th most abundant element on earth, averaging 0.034% of all Strontium output in 2005 igneous rock and is found chiefly as the form of the sulfate mineral celestite (SrSO4) and the carbonate strontianite (SrCO3). Of the two, celestite occurs much more frequently in sedimentary deposits of sufficient size to make development of mining facilities attractive. Strontianite would be the more useful of the two common minerals because strontium is used most often in the carbonate form, but few deposits have been discovered that are suitable for development.[5] The metal can be prepared by electrolysis of melted strontium chloride mixed with potassium chloride: Sr2+ + 2 e− → Sr 2 Cl− → Cl2 (g) + 2 e− Alternatively it is made by reducing strontium oxide with aluminium in a vacuum at a temperature at which strontium distills off. Three allotropes of the metal exist, with transition points at 235 and 540 °C. The largest commercially exploited deposits are found in England.
Applications As a pure metal strontium is used in strontium 90%-aluminium 10% alloys of an eutectic composition for the modification of aluminium-silicon casting alloys.[6] The primary use for strontium compounds is in glass for colour television cathode ray tubes to prevent X-ray emission.[7] [8] Other uses: •
89
•
90
Sr is the active ingredient in Metastron, a radiopharmaceutical used for bone pain secondary to metastatic bone cancer. The strontium acts like calcium and is preferentially incorporated into bone at sites of increased osteogenesis. This localization focuses the radiation exposure on the cancerous lesion. Sr has been used as a power source for radioisotope thermoelectric generators (RTGs). Sr produces about 0.93 watts of heat per gram (it is lower for the grade of 90Sr used in RTGs, which is strontium fluoride).[9] However, 90Sr has a lifetime approximately 3 times shorter and has a lower density than 238Pu, another RTG fuel. The main advantage of 90 Sr is that it is cheaper than 238Pu and is found in nuclear waste. 90
Strontium •
90
Sr is also used in cancer therapy. Its beta emission and long half-life is ideal for superficial radiotherapy. • Strontium is one of the constituents of AJ62 alloy, a durable magnesium alloy used in car and motorcycle engines by BMW. • Since Strontium is so similar to calcium, it is incorporated in the bone. All four isotopes are incorporated, in roughly similar proportions as they are found in nature (please see below). However the actual distribution of the isotopes tends to vary greatly from one geographical location to another. Thus analyzing the bone of an individual can help determine the region it came from. This approach helps to identify the ancient migration patterns as well as the origin of commingled human remains in battlefield burial sites. Strontium thus helps forensic scientists too. • Strontium is used in studies of neurotransmitter release in neurons. Like calcium, strontium facilitates synaptic vesicle fusion with the synaptic membrane. But unlike calcium, strontium causes asynchronous vesicle fusion. Therefore, replacing calcium in the culture medium with strontium allows scientists to measure the effects of a single vesicle fusion event, e.g., the size of the postsynaptic response elicited by the neurotransmitter content of a single vesicle.[10] [11] 87
Sr/86Sr ratios are commonly used to determine the likely provenance areas of sediment in natural systems, especially in marine and fluvial environments. Dasch (1969) showed that surface sediments of Atlantic displayed 87Sr/86Sr ratios that could be regarded as bulk averages of the 87Sr/86Sr ratios of geological terranes from adjacent landmasses.[12] A good example of a fluvial-marine system to which Sr isotope provenance studies have been successfully employed is the River Nile-Mediterranean system [13] [14] [15] . Due to the differing ages of the rocks that constitute the majority of the Blue and White Nile catchment areas of the changing provenance of sediment reaching the River Nile delta and East Mediterranean Sea can be discerned through Sr isotopic studies. Such changes are climatically controlled in the Late Quaternary. More recently, 87Sr/86Sr ratios have also been used to determine the source of ancient archaeological materials such as timbers and corn in Chaco Canyon, New Mexico[16] [17] . 87 Sr/86Sr ratios in teeth may also be used to track animal migrations [18] [19] or in criminal forensics. Strontium atoms are used in an experimental atomic clock with record-setting accuracy.[20]
Effect on the human body The human body absorbs strontium as if it were calcium. Due to the elements being sufficiently similar chemically, the stable forms of strontium might not pose a significant health threat -- in fact, the levels found naturally may actually be beneficial (see below) -but the radioactive 90Sr can lead to various bone disorders and diseases, including bone cancer. The strontium unit is used in measuring radioactivity from absorbed 90Sr. A recent in-vitro study conducted the NY College of Dental Sciences using strontium on osteoblasts showed marked improvement on bone-building osteoblasts.[21] An innovative drug made by combining strontium with ranelic acid has aided in bone growth, boosted bone density and lessened vertebral, peripheral and hip fractures.[22] [23] Women receiving the drug showed a 12.7% increase in bone density. Women receiving a placebo had a 1.6% decrease. Half the increase in bone density (measured by x-ray
6
Strontium densitometry) is attributed to the higher atomic weight of Sr compared with calcium, whereas the other half a true increase in bone mass. Strontium ranelate is registered as a prescription drug in Europe and many countries worldwide. It needs to be prescribed by a doctor, delivered by a pharmacist, and requires strict medical supervision. Currently (early 2007), it is not available in Canada or the United States. Several other salts of strontium such as strontium citrate or strontium carbonate are often presented as natural therapies and sold at a dose that is several hundred times higher than the usual strontium intake. Despite the lack of strontium deficit referenced in the medical literature and the lack of information about possible toxicity of strontium supplementation, such compounds can still be sold in the United States under the Dietary Supplements Health and Education Act of 1994. Their long-term safety and efficacy have never been evaluated on humans using large-scale medical trials.
References [1] P. Colarusso et al. (1996). " High-Resolution Infrared Emission Spectrum of Strontium Monofluoride (http:/ / bernath. uwaterloo. ca/ media/ 149. pdf)". J. Molecular Spectroscopy 175: 158. . [2] Murray, W.H. (1977). The Companion Guide to the West Highlands of Scotland. London: Collins. [3] " Strontian gets set for anniversary (http:/ / www. lochaber-news. co. uk/ news/ fullstory. php/ aid/ 2644/ Strontian_gets_set_for_anniversary. html)". Lochaber News. 19th June 2008. . [4] British Geological Survey (2009). World mineral production 2003–07 (http:/ / www. bgs. ac. uk/ mineralsuk/ downloads/ wmp_2003_2007. pdf). Keyworth, Nottingham: British Geological Survey. ISBN 978-0-85272-639-6. . Retrieved April 6, 2009. [5] Ober, Joyce A.. " Mineral Comodity Summaries 2008: Strontium (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ strontium/ mcs-2008-stron. pdf)" (PDF). United States Geological Survey. . Retrieved 2008-10-14. [6] " Aluminium – Silicon Alloys : Strontium Master Alloys for Fast Al-Si Alloy Modification from Metallurg Aluminium (http:/ / www. azom. com/ Details. asp?ArticleID=3353)". AZo Journal of Materials Online. . Retrieved 2008-10-14. [7] " Cathode Ray Tube Glass-To-Glass Recycling (http:/ / yosemite. epa. gov/ ee/ epa/ riafile. nsf/ vwAN/ S99-23. pdf)" (PDF). ICF Incorporated, USEP Agency. . Retrieved 2008-10-14. [8] Ober, Joyce A.; Polyak, Désirée E.. " Mineral Yearbook 2007: Strontium (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ strontium/ myb1-2007-stron. pdf)" (PDF). United States Geological Survey. . Retrieved 2008-10-14. [9] " What are the fuels for radioisotope thermoelectric generators? (http:/ / www. qrg. northwestern. edu/ projects/ vss/ docs/ Power/ 3-what-are-the-fuels-for-rtgs. html)". . [10] Miledi, R. (1966). "Strontium as a Substitute for Calcium in the Process of Transmitter Release at the Neuromuscular Junction". Nature 212: 1233. doi: 10.1038/2121233a0 (http:/ / dx. doi. org/ 10. 1038/ 2121233a0). [11] Hagler D.J., Jr, Goda Y. (2001). "Properties of synchronous and asynchronous release during pulse train depression in cultured hippocampal neurons". J. Neurophysiol. 85: 2324. [12] Dasch, J. (1969). "Strontium isotopes in weathering profiles, deep-sea sediments, and sedimentary rocks". Geochimica et Cosmochimica Acta 33 (12): 1521-1552. doi: 10.1016/0016-7037(69)90153-7 (http:/ / dx. doi. org/ 10. 1016/ 0016-7037(69)90153-7). [13] Krom, M. et al. (1999). "The characterisation of Saharan dusts and Nile particulate matter in surface sediments from the Levantine basin using Sr isotopes". Marine Geology 155 (3-4): 319-330. doi: 10.1016/S0025-3227(98)00130-3 (http:/ / dx. doi. org/ 10. 1016/ S0025-3227(98)00130-3). [14] Krom, M. D. et al. (2002). "Nile River sediment fluctuations over the past 7000 yr and their key role in sapropel development". Geology 30 (1): 71-74. doi: 10.1130/0091-7613(2002)030<0071:NRSFOT>2.0.CO;2 (http:/ / dx. doi. org/ 10. 1130/ 0091-7613(2002)030<0071:NRSFOT>2. 0. CO;2). [15] Talbot, M. R. et al. (2000). "Strontium isotope evidence for late Pleistocene reestablishment of an integrated Nile drainage network". Geology 28 (4): 343-346. doi: 10.1130/0091-7613(2000)28<343:SIEFLP>2.0.CO;2 (http:/ / dx. doi. org/ 10. 1130/ 0091-7613(2000)28<343:SIEFLP>2. 0. CO;2). [16] Benson, L., Cordell, L., Vincent, K., Taylor, H., Stein, J., Farmer, G., and Kiyoto, F. (2003). "Ancient maize from Chacoan great houses: where was it grown?". Proceedings of the National Academy of Sciences 100 (22):
7
Strontium 13111-13115. doi: 10.1073pnas.2135068100 (http:/ / dx. doi. org/ 10. 1073pnas. 2135068100). [17] English NB, Betancourt JL, Dean JS, Quade J. (2001). " Strontium isotopes reveal distant sources of architectural timber in Chaco Canyon, New Mexico (http:/ / www. pnas. org/ cgi/ pmidlookup?view=long& pmid=11572943)". Proc Natl Acad Sci U S A 98 (21): 11891-6. doi: 10.1073/pnas.211305498 (http:/ / dx. doi. org/ 10. 1073/ pnas. 211305498). PMID 11572943. . [18] Barnett-Johnson, Rachel (2007). "Identifying the contribution of wild and hatchery Chinook salmon (Oncorhynchus tshawytscha) to the ocean fishery using otolith microstructure as natural tags". Canadian Journal of Fisheries and Aquatic Sciences 64 (12): 1683-1692. doi: 10.1139/F07-129 (http:/ / dx. doi. org/ 10. 1139/ F07-129). [19] Porder, S., Paytan, A., and E.A. Hadly (2003). "Mapping the origin of faunal assemblages using strontium isotopes". Paleobiology 29 (2): 197-204. doi: 10.1666/0094-8373(2003)029<0197:MTOOFA>2.0.CO;2 (http:/ / dx. doi. org/ 10. 1666/ 0094-8373(2003)029<0197:MTOOFA>2. 0. CO;2). [20] . doi: 10.1126/science.1153341 (http:/ / dx. doi. org/ 10. 1126/ science. 1153341). [21] " The Effects of Strontium Citrate on Osteoblast Proliferation and Differentiation (http:/ / iadr. confex. com/ iadr/ 2007orleans/ techprogram/ abstract_89231. htm)". . Retrieved 2009-07-07. [22] Meunier PJ, Roux C, Seeman E et al. (2004). "effects of strontium ranelate on the risk of vertebral fracture in women with postmenopausal osteoporosis.". New England Journal of Medicine 350: 459–468. doi: 10.1056/NEJMoa022436 (http:/ / dx. doi. org/ 10. 1056/ NEJMoa022436). PMID 14749454. [23] Reginster JY, Seeman E, De Vernejoul MC et al. (2005). "Strontium ranelate reduces the risk of nonvertebral fractures in postmenopausal women with osteoporosis: treatment of peripheral osteoporosis (TROPOS) study". J Clin Metab. 90: 2816–2822. doi: 10.1210/jc.2004-1774 (http:/ / dx. doi. org/ 10. 1210/ jc. 2004-1774). PMID 15728210.
8
Article Sources and Contributors
Article Sources and Contributors Strontium Source: http://en.wikipedia.org/w/index.php?oldid=306651043 Contributors: 1337Garda, 65.68.87.xxx, Ad-blaster2007, Adamrush, Adm820, Ahoerstemeier, Alansohn, Anclation, AndonicO, Andres, AndriyK, Anil1956, Anthony, Anwar saadat, Archimerged, Arkuat, BGManofID, Basement12, Beetstra, Ben MacDui, Benbest, BerserkerBen, BillFlis, BlueEarth, Bobo192, BorgQueen, Boris Barowski, BrianGV, Bryan Derksen, CWii, CYD, Canthusus, Carnildo, ChicXulub, Chris Henniker, Chuchunezumi, Cmichael, Coralsites, Cureself, Curious1i, D, DV8 2XL, Da monster under your bed, Danielnez1, Darrien, David Latapie, Delta G, Derek Ross, Discospinster, Djinn112, DocCalmt, DocGonzo, DocNatural, DocNatural2007, Donarreiskoffer, Dougofborg, Download, DrBob, DrRiver, Drbones67, Dreadstar, DreamGuy, Duncan MacCall, Dwmyers, Dysepsion, EIFY, ERcheck, EamonnPKeane, Edgar181, El C, Emperorbma, Eoghan, Epolk, Erik Zachte, Eritain, Femto, FisherQueen, Folip2006, Fonzy, FredPoir, Frumpkin, Fvasconcellos, Gaius Cornelius, Gene Nygaard, Glen, Glorric, Greatpatton, Greg Allen, Grendelkhan, Hadal, Hak-kâ-ngìn, Hellbus, Hhs08, Hongooi, Hontogaichiban, IW.HG, Icairns, Icelight, ImperatorExercitus, Ixfd64, J.delanoy, Jacob.l345, Jamesy, Jaraalbe, Jeronimo, Jjb123, Jmb, Joanjoc, JodyB, John, Jose77, KPbIC, Kafziel, Kajasudhakarababu, Karada, Karl-Henner, Kinston eagle, Knowledge Seeker, Ksbrown, Kwamikagami, Kwksi, LAX, LOL, Larry_Sanger, Larrybobb, Levil, LoyalSoldier, LuigiManiac, Mais oui!, Marlith, Materialscientist, Mav, Mdf, Meekywiki, Mentifisto, Mgimpel, Mikegrant, Minesweeper, Myself0101, NawlinWiki, Nenglish, Nergaal, Nihiltres, Nk, Nrhenderson, Opelio, Phil Boswell, Philip Trueman, PierreAbbat, Pilotguy, Piperh, Plasmic Physics, Polonium, Poolkris, Pras, PseudoOne, Qxz, RTC, Razorflame, RedWolf, Remember, Rich Farmbrough, Riddley, Rjwilmsi, Robinpup, Romanm, Saperaud, SauliH, Savant13, Schneelocke, Sengkang, Shaddack, Shanew2, Skaffman, SkerHawx, Skunkboy74, Sl, Smallweed, Smokizzy, Sn0wflake, StaticGull, Stifynsemons, Stone, Suisui, Sunborn, Svante, TRUGROUP, Tagishsimon, Taral, Tetracube, That Guy, From That Show!, The Rambling Man, Theoneintraining, Thinghy, Thricecube, TigerShark, Tim Starling, Tlesher, Toatwilight, Tohd8BohaithuGh1, Tone, Vary, Versageek, Vivio Testarossa, Vsmith, Waggers, Walkerma, Watch37264, Werdan7, Werewolfking, Whosasking, Wiki.longa, Wilfred Glendon XXVI, Xaosflux, Xous, Yekrats, Yyy, Zack, 331 anonymous edits
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9
Yttrium
1
Yttrium 39
strontium ← yttrium → zirconium
Sc ↑
Y ↓
Lu Periodic Table - Extended Periodic Table
General Name, Symbol, Number
yttrium, Y, 39
Element category
transition metals
Group, Period, Block
3, 5, d
Appearance
silvery white
Standard atomic weight
88.90585(2) g·mol
Electron configuration
[Kr] 4d 5s
−1
1
Electrons per shell
2
2, 8, 18, 9, 2 Physical properties
Phase
solid
Density (near r.t.)
4.472 g·cm
Liquid density at m.p.
4.24 g·cm−3
Melting point
1799 K (1526 °C, 2779 °F)
Boiling point
3609 K (3336 °C, 6037 °F)
Heat of fusion
11.42 kJ·mol−1
Heat of vaporization
365 kJ·mol−1
Specific heat capacity
(25 °C) 26.53 J·mol−1·K−1
−3
Vapor pressure P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1883
2075
(2320)
(2627)
(3036)
(3607)
Atomic properties Crystal structure
hexagonal
Yttrium
2
Oxidation states
3, 2, 1
Electronegativity
1.22 (Pauling scale)
Ionization energies (more)
1st: 600 kJ·mol−1 2nd: 1180 kJ·mol−1 3rd: 1980 kJ·mol−1
Atomic radius
180 pm
Covalent radius
190±7 pm Miscellaneous [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(r.t.) (α, poly) 596 nΩ·m
Thermal conductivity
(300 K) 17.2 W·m
Thermal expansion
(r.t.) (α, poly) 10.6 µm/(m·K)
Speed of sound (thin rod)
(20 °C) 3300 m/s
Young's modulus
63.5 GPa
Shear modulus
25.6 GPa
Bulk modulus
41.2 GPa
−1
Poisson ratio
−1
·K
0.243
Brinell hardness
589 MPa
CAS registry number
7440-65-5 Most-stable isotopes
Main article: Isotopes of yttrium iso 87
Y
NA syn
half-life 3.35 d
DM ε γ
88
Y
89
Y
90
Y
91
Y
syn
100% syn
syn
106.6 d
DE (MeV) 0.48, 0.38D
ε
-
γ
1.83, 0.89
DP 87
Sr
88
Sr
-
89
Y is stable with 50 neutron
2.67 d
58.5 d
β−
2.28
90
γ
2.18
-
β−
1.54
91
γ
1.20
-
Zr
Zr
References
Yttrium (pronounced /ˈɪtriəm/) is a chemical element with symbol Y and atomic number 39. It is a silvery-metallic transition metal chemically similar to the lanthanoids and has historically been classified as a rare earth element.[2] Yttrium is almost always found
Yttrium combined with the lanthanoids in rare earth minerals and is never found in nature as a free element. Its only stable isotope, 89Y, is also its only naturally occurring isotope. In 1787, Carl Axel Arrhenius found a new mineral near Ytterby in Sweden and named it ytterbite, after the village. Johan Gadolin discovered yttrium's oxide in Arrhenius' sample in 1789,[3] and Anders Gustaf Ekeberg named the new oxide yttria. Elemental yttrium was first isolated in 1828 by Friedrich Wöhler.[4] The most important use of yttrium is in making phosphors, such as the red ones used in television cathode ray tube displays and in LEDs.[5] Other uses include the production of electrodes, electrolytes, electronic filters, lasers and superconductors; various medical applications; and as traces in various materials to enhance their properties. Yttrium has no known biological role, but exposure to yttrium compounds can cause lung disease in humans.[6]
Characteristics Properties Yttrium is a soft, silver-metallic, lustrous and highly crystalline transition metal in group 3. As expected by periodic trends, it is less electronegative than its predecessor in the group, scandium, more electronegative than its successor in the group, lanthanum, and less electronegative than the next member of period 5, zirconium.[7] [8] Yttrium is the first d-block element in the fifth period. The pure element is relatively stable in air in bulk form, due to passivation resulting from the formation of a protective oxide (Y2O3) film on its surface. This film can reach a thickness of 10 µm when yttrium is heated to 750 °C in water vapor.[9] When finely divided, however, yttrium is very unstable in air; shavings or turnings of the metal can ignite in air at temperatures exceeding 400 °C. [4] Yttrium nitride (YN) is formed when the metal is heated to 1000 °C in nitrogen.[9]
Similarity to the lanthanoids The similarities of yttrium to the lanthanoids are so strong that the element has historically been grouped with them as a rare earth element,[2] and is always found in nature together with them in rare earth minerals.[10] Chemically, yttrium resembles these elements more closely than its neighbor in the periodic table, scandium,[11] and if its physical properties were plotted against atomic number then it would have an apparent number of 64.5 to 67.5, placing it between the lanthanoids gadolinium and erbium.[12] It often also falls in the same range for reaction order,[9] resembling terbium and dysprosium at its chemical reactivity.[5] Yttrium is so close in size to the so-called 'Yttrium group' of heavy lanthanoid ions that in solution, it behaves as if it were one of them.[9] [13] Even though the lanthanoids are one row farther down the periodic table than yttrium, the similarity in atomic radius may be attributed to the lanthanoid contraction.[14] One of the few notable differences between the chemistry of yttrium and that of the lanthanoids is that yttrium is almost exclusively trivalent, whereas about half of the lanthanoids can have valences other than three.[9]
3
Yttrium
Compounds and reactions As a trivalent transition metal, yttrium forms various inorganic compounds, generally in the oxidation state of +3, by giving up all three of its valence electrons.[15] A good example is yttrium(III) oxide (Y2O3), also known as yttria, a six-coordinate white solid.[16] Yttrium forms a water-insoluble fluoride, hydroxide, and oxalate, but its bromide, chloride, iodide, nitrate and sulfate are all soluble in water.[9] The Y3+ ion is colorless in solution because of the absence of d and f electron shells.[9] Water readily reacts with yttrium and its compounds to form hydrogen gas and Y2O3.[10] Concentrated nitric and hydrofluoric acids do not rapidly attack yttrium, but other strong acids do.[9] With halogens, yttrium forms trihalides such as yttrium(III) fluoride (YF3), yttrium(III) chloride (YCl3), and yttrium(III) bromide (YBr3) at temperatures above roughly 200 °C.[6] Similarly, carbon, phosphorus, selenium, silicon and sulfur all form binary compounds with yttrium at elevated temperatures.[9] Organoyttrium chemistry is the study of compounds containing carbon–yttrium bonds. A few of these are known to have yttrium in the oxidation state 0.[17] [18] (The +2 state has been observed in chloride melts,[19] and +1 in oxide clusters in the gas phase.[20] ) Some trimerization reactions were observed by using organoyttrium compounds as catalysts.[18] These compounds use YCl3 as a starting material, which in turn is obtained from Y2O3 and concentrated hydrochloric acid and ammonium chloride.[21] [22] Hapticity is how a group of contiguous atoms of a ligand are coordinated to a central atom; it is indicated by the Greek character eta, η. Yttrium complexes were the first examples of complexes where carboranyl ligands were bound to a d0-metal center through a η7-hapticity.[18] Vaporization of the graphite intercalation compounds graphite–Y or graphite–Y2O3 leads to the formation of endohedral fullerenes such as Y@C82.[5] Electron spin resonance studies indicated the formation of Y3+ and (C82)3− ion pairs.[5] The carbides Y3C, Y2C, and YC2 can each hydrolyze to form hydrocarbons.[9]
4
Yttrium
Nucleosynthesis and isotopes Yttrium in the Solar System was created through stellar nucleosynthesis, mostly by the s-process (≈72%), but also by the r-process (≈28%).[23] The r-process consists of rapid neutron capture of lighter elements during supernova explosions. The s-process is a slow neutron capture of lighter elements inside pulsating red giant stars.[24] Yttrium isotopes are among the most common products of the nuclear fission of uranium occurring in nuclear explosions and nuclear reactors. In terms of waste management, the most important yttrium isotopes are 91 Y and 90Y, with half-lives of 58.51 days and 64 hours, respectively.[25] The first is formed directly from fission, while the latter, despite its short half-life, is in secular equilibrium with its long-lived parent isotope, strontium-90 (90Sr) with a half-life of 29 years.[4] All group 3 elements have an odd number of protons and therefore have few stable isotopes.[7] Yttrium itself Mira is an example of the type of red has only one stable isotope, 89Y, which is also its only 89 giant star where most of the yttrium in naturally occurring one. Y is thought to be more the solar system was created. abundant than it otherwise would be, due in part to the s-process which allows enough time for isotopes created by other processes to decay by electron emission (neutron → proton).[24] [26] Such a slow process tends to favor isotopes with mass numbers (A = protons + neutrons) around 90, 138 and 208, which have unusually stable atomic nuclei with 50, 82 and 126 neutrons, respectively.[24] [27] [4] 89Y has a mass number close to 90 and has 50 neutrons in its nucleus. At least 32 synthetic isotopes of yttrium have been observed, ranging in mass number from 76 to 108.[25] The least stable of these is 106Y with a half-life of >150 ns (76Y has a half-life of >200 ns) and the most stable is 88Y with a half-life of 106.626 days.[25] Besides the isotopes 91Y, 87Y, and 90Y, with half lives of 58.51 days, 79.8 hours, and 64 hours, respectively, all the other isotopes have half lives of less than a day and most of those have half-lives of less than an hour.[25] Yttrium isotopes with mass numbers at or below 88 decay primarily by positron emission (proton → neutron) to form strontium (Z = 38) isotopes.[25] Yttrium isotopes with mass numbers at or above 90 decay primarily by electron emission (neutron → proton) to form zirconium (Z = 40) isotopes.[25] Isotopes with mass numbers at or above 97 are also known to have minor decay paths of β− delayed neutron emission.[28] Yttrium has at least 20 metastable or excited isomers ranging in mass number from 78 to 102.[25] [29] Multiple excitation states have been observed for 80Y and 97Y.[25] While most of yttrium's isomers are expected to be less stable than their ground state, 78mY, 84mY, 85mY, 96m Y, 98m1Y, 100mY, and 102mY have longer half-lives than their ground states, as these isomers decay by beta decay rather than isomeric transition.[28]
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History In 1787, army lieutenant and part-time chemist Carl Axel Arrhenius found a heavy black rock in an old quarry near the Swedish village of Ytterby (now part of the Stockholm Archipelago).[3] Thinking that it was an unknown mineral containing the newly discovered element tungsten,[30] he named it ytterbite[31] and sent samples to various chemists for further analysis.[3] Johan Gadolin at the University of Åbo identified a new oxide or "earth" in Arrhenius' sample in 1789, and published his completed analysis in 1794.[32] [33] Anders Gustaf Ekeberg confirmed this in 1797 and named the new oxide yttria.[34] In the decades after Antoine Lavoisier developed the first modern definition of chemical elements, it was believed that earths could be reduced to their elements, meaning that the discovery of a new earth was equivalent to the discovery of the element within, which in this case would have been yttrium.[35] In 1843, Carl Gustav Mosander found that samples of yttria Johan Gadolin discovered actually contained three oxides: white yttrium oxide (yttria), yttrium oxide. yellow terbium oxide (confusingly, this was called 'erbia' at the time) and rose-colored erbium oxide (called 'terbia' at the time).[36] A fourth oxide, ytterbium oxide, was isolated in 1878 by Jean Charles Galissard de Marignac.[37] New elements would later be isolated from each of those oxides, and each element was named, in some fashion, after Ytterby, the village near the quarry in which they were found (see ytterbium, terbium, and erbium).[38] In the following decades, seven other new metals were discovered in "Gadolin's yttria".[3] Since yttria was a mineral after all and not an oxide, Martin Heinrich Klaproth renamed it gadolinite in honor of Gadolin.[3] Yttrium metal was first isolated in 1828 when Friedrich Wöhler heated anhydrous yttrium(III) chloride with potassium:[39] [40] YCl3 + 3 K → 3 KCl + Y Until the early 1920s, the chemical symbol Yt was used for the element, after which Y came into common use.[41] In 1987, yttrium barium copper oxide was found to achieve high-temperature superconductivity.[42] It was only the second material known to exhibit this property,[42] and it was the first known material to achieve superconductivity above the (economically important) boiling point of nitrogen.[43]
Yttrium
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Occurrence Abundance Yttrium is found in most rare earth minerals,[8] as well as some uranium ores, but is never found in nature as a free element.[44] About 31 ppm of the Earth's crust is yttrium,[5] making it the 28th most abundant element there, and 400 times more common than silver.[45] Yttrium is found in soil in concentrations between 10 and 150 ppm (dry weight average of 23 ppm) and in sea water at 9 ppt.[45] Lunar rock samples collected during the Apollo program have a relatively high yttrium content.[38] Xenotime crystals contain yttrium.
Yttrium has no known biological role, though it is
found in most, if not all, organisms and tends to concentrate in the liver, kidney, spleen, lungs, and bones of humans.[46] There is normally as little as 0.5 milligrams found within the entire human body; human breast milk contains 4 ppm.[47] Yttrium can be found in edible plants in concentrations between 20 ppm and 100 ppm (fresh weight), with cabbage having the largest amount.[47] With up to 700 ppm, the seeds of woody plants have the highest known concentrations.[47]
Production The chemical similarity of yttrium with the lanthanoids leads it to being enriched by the same processes and ends up in ores containing lanthanoids, forming rare earth minerals. A slight separation is recognized between the light (LREE) and the heavy rare earth elements (HREE) but this separation is never complete. Yttrium is concentrated in the HREE group even though it has a lower atomic mass.[48] [49]
Yttrium
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There are four main sources for REEs:[50] • Carbonate and fluoride containing ores such as the LREE bastnäsite ([(Ce, La, etc.)(CO3)F]) contain an average of 0.1%[4] [48] of yttrium compared to the 99.9% for the 16 other REEs.[48] The main source for bastnäsite from the 1960s to the 1990s was the Mountain Pass mine in California, making the United States the largest producer of REEs.[48] [50] • Monazite ([(Ce, La, etc.)PO4]), which is mostly phosphate, is a placer deposit of sand that is created by the transportation and gravitational separation of eroded granite. Monazite as a LREE ore contains 2%[48] (or 3%)[51] of yttrium. The largest deposits were found in India and Brazil in the early 19th century, making these two countries the largest producers of yttrium in the first half of that century.[48] [50]
A piece of yttrium. Yttrium is difficult to separate from other rare earth elements.
• Xenotime, a REE phosphate, is the main HREE ore containing up to 60% of yttrium as yttrium phosphate (YPO4).[48] The largest mine for this mineral is the Bayan Obo deposit in China, making China the largest exporter for HREE since the closure of the Mountain Pass mine in the 1990s.[48] [50] • Ion absorption clays or Lognan clays are the weathering products of granite and contain only 1% of REEs.[48] The final ore concentrate can contain up to 8% of yttrium. Ion absorption clays are mostly mined in southern China.[48] [50] [52] Yttrium is also found in samarskite and fergusonite.[45] It is difficult to separate yttrium from other rare earths. One method to obtain pure yttrium from the mixed oxide ores is to dissolve the oxide in sulfuric acid and fractionate it by ion exchange chromatography. With the addition of oxalic acid, the yttrium oxalate precipitates. The oxalate is converted into the oxide by heating under oxygen. By reacting the resulting yttrium oxide with hydrogen fluoride, yttrium fluoride is obtained.[53] Annual world production of yttrium oxide had reached 600 tonnes by 2001, with reserves estimated at 9 million tonnes.[45] Only a few tonnes of yttrium metal are produced each year by reducing yttrium fluoride to a metal sponge with calcium magnesium alloy. The temperature of an arc furnace of above 1,600 °C is sufficient to melt the yttrium.[45] [53]
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Applications Consumer Yttria (Y2O3) can serve as host lattice for doping with Eu3+ cations as well as reactant to gain doped yttrium orthovanadate YVO4:Eu3+ or yttrium oxide sulfide Y2O2S:Eu3+ phosphors that give the red color in color television picture tubes,[4] [5] [54] though the red color itself is actually emitted from the europium while the yttrium collects energy from the electron gun and passes it to the phosphor.[55] Yttrium compounds can serve as host lattices for doping with different lanthanoid cations. Besides Eu3+ also Tb3+ can be used as a doping agent leading to green luminescence. Yttria is also used as a sintering additive in the production of porous silicon nitride[56] and as a common starting material for both material science and for producing other compounds of yttrium.
Yttrium is one of the elements used to make the red color in CRT televisions.
Yttrium compounds are used as a catalyst for ethylene polymerization.[4] As a metal, it is used on the electrodes of some high-performance spark plugs.[57] Yttrium is also used in the manufacturing of gas mantles for propane lanterns as a replacement for thorium, which is radioactive.[58] Developing uses include yttrium-stabilized zirconia in particular as a solid electrolyte and as an oxygen sensor in automobile exhaust systems.[5]
Garnets Yttrium is used in the production of a large variety of synthetic garnets,[59] and yttria is used to make yttrium iron garnets (YIG), which are very effective microwave filters.[4] Yttrium, iron, aluminium, and gadolinium garnets (e.g. Y3Fe5O12 and Y3Al5O12) have important magnetic properties.[4] YIG is also very efficient as an acoustic energy transmitter and transducer.[60] Yttrium aluminium garnet (Y3Al5O12 or YAG) has a hardness of 8.5 and is also used as a gemstone in jewelry (simulated diamond).[4] Cerium-doped yttrium aluminium garnet (YAG:Ce) crystals are used as phosphors to make white LEDs.[61] [62] [63]
YAG, yttria, yttrium lithium fluoride (LiYF4), and yttrium orthovanadate (YVO4) are used in combination with dopants such as neodymium, erbium, ytterbium in near-infrared lasers.[64] [65] YAG lasers have the ability to operate at high power and are used for drilling into and cutting metal.[51] The single crystals of doped YAG are normally produced by the Czochralski process.[66]
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Material enhancer Small amounts of yttrium (0.1 to 0.2%) have been used to reduce the grain sizes of chromium, molybdenum, titanium, and zirconium.[67] It is also used to increase the strength of aluminium and magnesium alloys.[4] The addition of yttrium to alloys generally improves workability, adds resistance to high-temperature recrystallization and significantly enhances resistance to high-temperature oxidation (see graphite nodule discussion below).[55] Yttrium can be used to deoxidize vanadium and other non-ferrous metals.[4] Yttria is used to stabilize the cubic form of zirconia for use in jewelry.[68] Yttrium has been studied for possible use as a nodulizer in the making of nodular cast iron which has increased ductility (the graphite forms compact nodules instead of flakes to form nodular cast iron).[4] Yttrium oxide can also be used in ceramic and glass formulas, since it has a high melting point and imparts shock resistance and low thermal expansion characteristics.[4] It is therefore used in camera lenses.[45]
Medical The radioactive isotope yttrium-90 is used in drugs such as Yttrium Y 90-DOTA-tyr3-octreotide and Yttrium Y 90 ibritumomab tiuxetan for the treatment of various cancers, including lymphoma, leukemia, ovarian, colorectal, pancreatic, and bone cancers.[47] It works by adhering to monoclonal antibodies, which in turn bind to cancer cells and kill them via intense β-radiation from the yttrium-90 (see Monoclonal antibody therapy).[69] Needles made of yttrium-90, which can cut more precisely than scalpels, have been used to sever pain-transmitting nerves in the spinal cord,[30] and yttrium-90 is also used to carry out radionuclide synovectomy in the treatment of inflamed joints, especially knees, in sufferers of conditions such as rheumatoid arthritis.[70] A neodymium-doped yttrium-aluminium-garnet laser has been used in an experimental, robot-assisted radical prostatectomy in canines in an attempt to reduce collateral nerve and tissue damage,[71] whilst the erbium-doped ones are starting to be used in cosmetic skin resurfacing.[5]
Superconductors Yttrium was used in the yttrium barium copper oxide (YBa2Cu3O7, aka 'YBCO' or '1-2-3') superconductor developed at the University of Alabama and the University of Houston in 1987.[42] This superconductor operated at 93 K, notable because this is above liquid nitrogen's boiling point (77.1 K).[42] As the price of liquid nitrogen is lower than that of liquid helium, which has to be used for the metallic superconductors, the operating costs would decrease. YBCO superconductor
The actual superconducting material is often written as YBa2Cu3O7−d, where d must be less than 0.7 if the
Yttrium material is to be superconducting. The reason for this is still not clear, but it is known that the vacancies occur only in certain places in the crystal, the copper oxide planes and chains, giving rise to a peculiar oxidation state of the copper atoms, which somehow leads to the superconducting behaviour. The theory of low temperature superconductivity has been well understood since the so-called BCS theory was put forward in 1957. It is based on a peculiarity of the interaction between 2 electrons in a crystal lattice. However, BCS theory does not explain high temperature superconductivity, and its precise mechanism is still a mystery. What is known is that the composition of the copper-oxide materials has to be precisely controlled if superconductivity is to occur.[72] The created material was a black and green, multi-crystal, multi-phase mineral. Researchers are studying a class of materials known as perovskites that are alternative mixtures of these elements, hoping to eventually develop a practical high-temperature superconductor.[51]
Precautions Water soluble compounds of yttrium are considered mildly toxic, while its insoluble compounds are non-toxic.[47] In experiments on animals, yttrium and its compounds caused lung and liver damage, though toxicity varies with different yttrium compounds. In rats, inhalation of yttrium citrate caused pulmonary edema and dyspnea, while inhalation of yttrium chloride caused liver edema, pleural effusions, and pulmonary hyperemia.[6] Exposure to yttrium compounds in humans may cause lung disease.[6] Workers exposed to airborne yttrium europium vanadate dust experienced mild eye, skin, and upper respiratory tract irritation—though this may have been caused by the vanadium content rather than the yttrium.[6] Acute exposure to yttrium compounds can cause shortness of breath, coughing, chest pain, and cyanosis.[6] NIOSH recommends a time-weighted average limit of 1 mg/m3 and an IDLH of 500 mg/m3.[73] Yttrium dust is flammable.[6]
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] IUPAC 2005 [3] Van der Krogt 2005 [4] CRC 2008, v.4, p. 41 [5] Cotton 2008 [6] OSHA 2007 [7] Greenwood 1997, p. 946 [8] Hammond [9] Daane 1968, p. 817 [10] Emsley 2001, p. 498 [11] Daane 1968, p. 810 [12] Daane 1968, p. 815 [13] Greenwood 1997, p. 945 [14] Greenwood 1997, p. 1234 [15] Greenwood 1997, p. 948 [16] Greenwood 1997, p. 947 [17] Cloke 1993 [18] Schumann 2006 [19] Mikheev 1992 [20] Kang 2005
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[21] Turner 1920, p. 492 [22] Spencer 1919, p. 135 [23] Pack 2007 [24] Greenwood 1997, pp. 12–13 [25] NNDC 2008 [26] Essentially, a neutron becomes a proton while an electron and antineutrino are emitted. [27] This stability is thought to result from very low neutron cross-sections (Greenwood 1997, pp. 12—13). Electron emission of isotopes with those mass numbers is simply less prevalent due to this stability, resulting in them having a higher abundance. [28] Audi 2003 [29] Metastable isomers have higher-than-normal energy states than the corresponding non-excited nucleus and these states last until a gamma ray or conversion electron is emitted from the isomer. They are designated by an 'm' being placed next to the isotope's mass number. [30] Emsley 2001, p. 496 [31] Ytterbite was named after the village it was discovered near, plus the -ite ending to indicate it was a mineral. [32] Gadolin 1794 [33] Stwertka 1998, p. 115 says that the identification occurred in 1789 but is silent on when the announcement was made. Van der Krogt 2005 cites the original publication, with the year 1794, by Gadolin. [34] [35] [36] [37]
Greenwood 1997, p. 944 Earths were given an -a ending and new elements are normally given an -ium ending Mosander 1843 Britannica 2005, "ytterbium"
[38] Stwertka 1998, p. 115 [39] Heiserman 1992, p. 150 [40] Wöhler 1828 [41] Coplen and Peiser 1998 [42] Wu et al.. 1987 [43] Tc for YBCO is 93 K and the boiling point of nitrogen is 77 K. [44] Lenntech, "yttrium" [45] Emsley 2001, p. 497 [46] Mac Donald et al.. 1952 [47] Emsley 2001, p. 495 [48] Morteani 1991 [49] Kanazawaa 2006 [50] Naumov 2008 [51] Stwertka 1998, p. 116 [52] Zuoping 1996 [53] Holleman 1985 [54] Emsley 2001, p. 497 says that "Yttrium oxysulfide, doped with europium (III), is used as the standard red component in colour televisions". [55] [56] [57] [58] [59] [60] [61] [62] [63] [64] [65] [66] [67] [68] [69] [70]
Daane 1968, p. 818 US patent 5935888 Carley 2000 Addison 1985 Jaffe 1951 Hosseinivajargah 2007 US patent 6409938 GIA 1995 Kiss and Pressley 1996 Kong et al. 2005 Tokurakawa et al.. 2007 Aleksandar et al.. 2002 PIDC contributors Berg 2002 Adams et al.. 2004 Fischer 2002
[71] Gianduzzo 2008 [72] Imperial College (http:/ / www. ch. ic. ac. uk/ rzepa/ mim/ century/ html/ ybco_text. htm) [73] NIOSH 2005
Yttrium
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Yttrium • GIA contributors (1995). GIA Gem Reference Guide. Gemological Institute of America. ISBN 0-87311-019-6. • Gianduzzo, Troy; Colombo Jr, Jose R.; Haber, Georges-Pascal; Hafron, Jason; Magi-Galluzzi, Cristina; Aron, Monish; Gill, Inderbir S.; Kaouk, Jihad H. (September 2008). "Laser robotically assisted nerve-sparing radical prostatectomy: a pilot study of technical feasibility in the canine model". BJU International (Cleveland: Glickman Urological Institute) 102 (5): 598. doi: 10.1111/j.1464-410X.2008.07708.x (http:/ / dx. doi. org/ 10. 1111/ j. 1464-410X. 2008. 07708. x). • Golubović, Aleksandar V.; Nikolić, Slobodanka N.; Gajić, Radoš; Đurić, Stevan; Valčić, Andreja (2002). "The growth of Nd: YAG single crystals". Journal of the Serbian Chemical Society 67 (4): 91–300. doi: 10.2298/JSC0204291G (http:/ / dx. doi. org/ 10. 2298/ JSC0204291G). • Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. ISBN 0-7506-3365-4. • Hammond, C. R.. " Yttrium (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elements. pdf)" (pdf). The Elements. Fermi National Accelerator Laboratory. pp. 4–33. http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elements. pdf. Retrieved 2008-08-26. • Heiserman, David L. (1992). "Element 39: Yttrium". Exploring Chemical Elements and their Compounds. New York: TAB Books. pp. 150–152. ISBN 0-8306-3018-X. • Holleman, Arnold F.; Egon Wiberg, Nils Wiberg (1985). Lehrbuch der Anorganischen Chemie (91–100 ed.). Walter de Gruyter. pp. 1056–1057. ISBN 3-11-007511-3. • Vajargah, S. Hosseini (2007). "Preparation and characterization of yttrium iron garnet (YIG) nanocrystalline powders by auto-combustion of nitrate-citrate gel". Journal of Alloys and Compounds 430 (1–2): 339–343. doi: 10.1016/j.jallcom.2006.05.023 (http:/ / dx. doi. org/ 10. 1016/ j. jallcom. 2006. 05. 023). • IUPAC contributors (2005). Edited by N G Connelly and T Damhus (with R M Hartshorn and A T Hutton). ed (PDF). Nomenclature of Inorganic Chemistry: IUPAC Recommendations 2005 (http:/ / www. iupac. org/ publications/ books/ rbook/ Red_Book_2005. pdf). RSC Publishing. pp. 51. ISBN 0-85404-438-8. http:/ / www. iupac. org/ publications/ books/ rbook/ Red_Book_2005. pdf. Retrieved 2007-12-17. • Jaffe, H.W. (1951). " The role of yttrium and other minor elements in the garnet group (http:/ / www. minsocam. org/ ammin/ AM36/ AM36_133. pdf)" (pdf). American Mineralogist: 133–155. http:/ / www. minsocam. org/ ammin/ AM36/ AM36_133. pdf. Retrieved 2008-08-26. • Kanazawaa, Yasuo; Masaharu Kamitani (2006). "Rare earth minerals and resources in the world". Journal of Alloys and Compounds 408–412: 1339–1343. doi: 10.1016/j.jallcom.2005.04.033 (http:/ / dx. doi. org/ 10. 1016/ j. jallcom. 2005. 04. 033). • Kang, Weekyung; E. R. Bernstein (2005). " Formation of Yttrium Oxide Clusters Using Pulsed Laser Vaporization (http:/ / newjournal. kcsnet. or. kr/ main/ j_search/ j_download. htm?code=B050237)". Bull. Korean Chem. Soc. 26 (2): 345–348. http:/ / newjournal. kcsnet. or. kr/ main/ j_search/ j_download. htm?code=B050237. • Kiss, Z. J.; Pressley, R. J. (October 1966). " Crystalline solid lasers (http:/ / ieeexplore. ieee. org/ xpls/ abs_all. jsp?arnumber=1447042)". Proceedings of the IEEE. 54. IEEE. pp. 1236–1248. issn: 0018-9219. http:/ / ieeexplore. ieee. org/ xpls/ abs_all. jsp?arnumber=1447042. Retrieved 2008-08-16.
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Yttrium • Kong, J.; Tang, D. Y.; Zhao, B.; Lu, J.; Ueda, K.; Yagi, H. and Yanagitani, T. (2005). "9.2-W diode-pumped Yb:Y2O3 ceramic laser". Applied Physics Letters 86: 116. doi: 10.1063/1.1914958 (http:/ / dx. doi. org/ 10. 1063/ 1. 1914958). • Lenntech contributors. " yttrium (http:/ / www. lenntech. com/ periodic-chart-elements/ y-en. htm)". Lenntech. http:/ / www. lenntech. com/ periodic-chart-elements/ y-en. htm. Retrieved 2008-08-26. • MacDonald, N. S.; R. E. Nusbaum, G. V. Alexander (1952). " The Skeletal Deposition of Yttrium (http:/ / www. jbc. org/ cgi/ reprint/ 195/ 2/ 837. pdf)" (PDF). Journal of Biological Chemistry 195: 837–841. http:/ / www. jbc. org/ cgi/ reprint/ 195/ 2/ 837. pdf. • Nikolai B., Mikheev (1992). "The anomalous stabilisation of the oxidation state 2+ of lanthanides and actinides". Russian Chemical Reviews 61 (10): 990–998. doi: 10.1070/RC1992v061n10ABEH001011 (http:/ / dx. doi. org/ 10. 1070/ RC1992v061n10ABEH001011). • Morteani, Giulio (01 Aug 1991). " The rare earths; their minerals, production and technical use (http:/ / eurjmin. geoscienceworld. org/ cgi/ content/ abstract/ 3/ 4/ 641)". European Journal of Mineralogy; August; v.; no.; p. 3 (4): 641–650. http:/ / eurjmin. geoscienceworld. org/ cgi/ content/ abstract/ 3/ 4/ 641. • Carl Gustav, Mosander (1843). "Ueber die das Cerium begleitenden neuen Metalle Lathanium und Didymium, so wie über die mit der Yttererde vorkommen-den neuen Metalle Erbium und Terbium" (in German). Annalen der Physik und Chemie 60 (2): 297–315. doi: 10.1002/andp.18431361008 (http:/ / dx. doi. org/ 10. 1002/ andp. 18431361008). • Naumov, A. V. (2008). " Review of the World Market of Rare-Earth Metals (http:/ / www. springerlink. com/ content/ y8925j378w4u4175/ )". Russian Journal of Non-Ferrous Metals 49 (1): 14–22. http:/ / www. springerlink. com/ content/ y8925j378w4u4175/ . • NIOSH contributors (September 2005). " Yttrium (http:/ / www. cdc. gov/ niosh/ npg/ npgd0673. html)". NIOSH Pocket Guide to Chemical Hazards. National Institute for Occupational Safety and Health. http:/ / www. cdc. gov/ niosh/ npg/ npgd0673. html. Retrieved 2008-08-03. • NNDC contributors (2008). " Chart of Nuclides (http:/ / www. nndc. bnl. gov/ chart/ )". in Alejandro A. Sonzogni (Database Manager). Upton, New York: National Nuclear Data Center, Brookhaven National Laboratory. http:/ / www. nndc. bnl. gov/ chart/ . Retrieved 2008-09-13. • OSHA contributors (2007-01-11). " Occupational Safety and Health Guideline for Yttrium and Compounds (http:/ / www. osha. gov/ SLTC/ healthguidelines/ yttriumandcompounds/ recognition. html)". United States Occupational Safety and Health Administration. http:/ / www. osha. gov/ SLTC/ healthguidelines/ yttriumandcompounds/ recognition. html. Retrieved 2008-08-03. (public domain text) • Pack, Andreas; Sara S. Russell, J. Michael G. Shelley and Mark van Zuilen (2007). "Geoand cosmochemistry of the twin elements yttrium and holmium". Geochimica et Cosmochimica Acta 71 (18): 4592–4608. doi: 10.1016/j.gca.2007.07.010 (http:/ / dx. doi. org/ 10. 1016/ j. gca. 2007. 07. 010). • PIDC contributors. Rare Earth metals & compounds (http:/ / www. pidc. com/ products_imaterials_oth. html). Pacific Industrial Development Corporation. http:/ / www. pidc. com/ products_imaterials_oth. html. Retrieved 2008-08-26. • Schumann, Herbert; Fedushkin, Igor L. (2006). "Scandium, Yttrium & The Lanthanides: Organometallic Chemistry". Encyclopedia of Inorganic Chemistry. doi:
15
Yttrium 10.1002/0470862106.ia212 (http:/ / dx. doi. org/ 10. 1002/ 0470862106. ia212). • Spencer, James F. (1919). The Metals of the Rare Earths (http:/ / books. google. com/ books?id=W2zxN_FLQm8C& pg=PA135& dq="Yttrium+ chloride"& lr=& as_brr=1). New York: Longmans, Green, and Co. pp. 135. http:/ / books. google. com/ books?id=W2zxN_FLQm8C& pg=PA135& dq=%22Yttrium+ chloride%22& lr=& as_brr=1. Retrieved 2008-08-12. • Stwertka, Albert (1998). "Yttrium". Guide to the Elements (Revised ed.). Oxford University Press. pp. 115–116. ISBN 0-19-508083-1. • Tokurakawa, M.; Takaichi, K.; Shirakawa, A.; Ueda, K.; Yagi, H.; Yanagitani, T. and Kaminskii, A. A. (2007). "Diode-pumped 188 fs mode-locked Yb3+:Y2O3 ceramic laser". Applied Physics Letters 90: 071101. doi: 10.1063/1.2476385 (http:/ / dx. doi. org/ 10. 1063/ 1. 2476385). • Turner, Jr., Francis M.; Berolzheimer, Daniel D.; Cutter, William P.; Helfrich, John (1920). The Condensed Chemical Dictionary (http:/ / books. google. com/ books?id=y8y0XE0nsYEC& pg=PA492& dq="Yttrium+ chloride"& lr=& as_brr=1). New York: Chemical Catalog Company. pp. 492. http:/ / books. google. com/ books?id=y8y0XE0nsYEC& pg=PA492& dq=%22Yttrium+ chloride%22& lr=& as_brr=1. Retrieved 2008-08-12. • US patent 4533317 (http:/ / v3. espacenet. com/ textdoc?DB=EPODOC& IDX=US4533317), "Yttrium oxide mantles for fuel-burning lanterns", granted 1985-08-06 , assigned to The Coleman Company, Inc. • US patent 5734166 (http:/ / v3. espacenet. com/ textdoc?DB=EPODOC& IDX=US5734166), "Low-energy neutron detector based upon lithium lanthanide borate scintillators", granted 1998-03-31 , assigned to Mission Support Inc • US patent 5935888 (http:/ / v3. espacenet. com/ textdoc?DB=EPODOC& IDX=US5935888), "Porous silicon nitride with rodlike grains oriented", granted 1999-08-10 , assigned to Agency Ind Science Techn (JP) and Fine Ceramics Research Ass (JP),. • US patent 6409938 (http:/ / v3. espacenet. com/ textdoc?DB=EPODOC& IDX=US6409938), "Aluminum fluoride flux synthesis method for producing cerium doped YAG", granted 2002-06-25 , assigned to General Electrics • van der Krogt, Peter (2005-05-05). " 39 Yttrium (http:/ / elements. vanderkrogt. net/ elem/ y. html)". Elementymology & Elements Multidict. http:/ / elements. vanderkrogt. net/ elem/ y. html. Retrieved 2008-08-06. • Wöhler, Friedrich (1828). "Ueber das Beryllium und Yttrium". Annalen der Physik 89 (8): 577–582. doi: 10.1002/andp.18280890805 (http:/ / dx. doi. org/ 10. 1002/ andp. 18280890805). • Wu, M. K.; Ashburn, J. R.; Torng, C. J.; Hor, P. H.; Meng, R. L.; Gao, L.; Huang, Z. J.; Wang, Y. Q. and Chu, C. W. (1987). "Superconductivity at 93 K in a New Mixed-Phase Y-Ba-Cu-O Compound System at Ambient Pressure". Physical Review Letters 58: 908–910. doi: 10.1103/PhysRevLett.58.908 (http:/ / dx. doi. org/ 10. 1103/ PhysRevLett. 58. 908). • Zheng, Zuoping; Lin Chuanxian (1996). "The behaviour of rare-earth elements (REE) during weathering of granites in southern Guangxi, China". Chinese Journal of Geochemistry 15 (4): 344–352. doi: 10.1007/BF02867008 (http:/ / dx. doi. org/ 10. 1007/ BF02867008).
16
Yttrium
External links • WebElements.com – Yttrium (http:/ / www. webelements. com/ webelements/ elements/ text/ Y/ index. html) pnb:میرتیا
17
Article Sources and Contributors
Article Sources and Contributors Yttrium Source: http://en.wikipedia.org/w/index.php?oldid=308250152 Contributors: AKAF, Ahoerstemeier, Akamad, Alan.duffield, Alansohn, Alexhuddart, Andres, Andrew Kanode, AnnaFrance, Antandrus, Archimerged, Arkuat, Art LaPella, Ayecee, Baccyak4H, Barnaby Ryans, Beetstra, Bender235, BlueEarth, Bobby.zooomer, Bobo192, Bryan Derksen, Burzmali, C.A.T.S. CEO, C.Fred, CStyle, CYD, Cabman, Cacahueten, Californium-256, Capricorn42, Carnildo, Christian List, Computer boy, Conversion script, Cosmic Latte, Cryptic C62, Crystal whacker, Dale101usa, Dalliance, Darrien, David Latapie, Deglr6328, Delta G, Dengero, Derek.cashman, Dfnj123, Dominick, Domitori, Donarreiskoffer, Doulos Christos, Dread Specter, Dwmyers, DÅ‚ugosz, E F F E C T, EchetusXe, Ecrowell4719, Elysdir, Emperorbma, Enviroboy, Eog1916, Epbr123, Erik Zachte, Erik9, Erkcan, Error4567890, Evlekis, Eweisser, Faradayplank, Femto, Freddyd945, GDonato, Galoubet, Ggonnell, Giftlite, Goodtimber, Graibeard, Grendelkhan, Gökhan, Haemo, Haham hanuka, Hak-kâ-ngìn, Hankwang, HappyCamper, Harrymph, Heimstern, Heterodoxphilomath, Hracjk, IanOsgood, Icairns, Ideyal, Indiealtphreak, Iridescent, Itub, J.delanoy, JForget, Jake Wartenberg, Jaraalbe, Jaxl, Jeff G., Jessi1989, Jkasd, Jkl, Joanjoc, John, JohnyDog, Jorourke92, Jose77, Jpk, Karl-Henner, Kenmcl2, Kerttie, King of Hearts, Kingpin13, Ktsquare, Kurykh, Kwamikagami, L337p4wn, LA2, LMB, LSU260, Lazylaces, LeoNomis, Lewis R, Lexicon, Ling.Nut, Livajo, Lord asriel, LovesMacs, MBisanz, MER-C, MKoltnow, Maddie!, Maralia, Martin451, Master of Puppets, Materialscientist, Mav, Mgimpel, Mic, Michael Devore, Mikhail Klassen, Minesweeper, Miranda, Mkweise, Mmccalpin, Mortdefides, Mouser, Mr0t1633, Mygerardromance, Ndsg, NeilN, Nergaal, Nihiltres, Octoferret, Onevalefan, Pastel kitten, Paulc206, Pcb21, PeteThePill, Petergans, Pevernagie, Plasmic Physics, Plexust, Poolkris, Poor Yorick, Pras, Prolinol, Puchiko, RTC, RainbowOfLight, RandomP, Redtitan, Redvers, Remember, Reza kalani, Rhanyeia, Rich Farmbrough, Rjstott, Rjwilmsi, RobertG, Romanm, Romeo400, Rrburke, Rursus, Salsa Shark, SandyGeorgia, Saperaud, Schneelocke, Sengkang, Shaddack, Shanel, Shimgray, Shirulashem, Sillyfolkboy, Silver Edge, Sinneed, Sionus, Sjc196, Sjö, Skippiikai, Sl, Solid State, Spaldinggrammerboy, Squash, Stan J Klimas, Steve Hart, Stifynsemons, Stone, Suisui, Sunborn, T-rex, Tagishsimon, Tetracube, Thepuppetmasterofspam, Thricecube, Titoxd, Tony1, Tsemii, Ulric1313, Vicki Rosenzweig, Violentbob, Vsmith, Warut, WaysToEscape, WereSpielChequers, Willking1979, Yekrats, Yyy, Zany Zebu Redivivus, 282 anonymous edits
Image Sources, Licenses and Contributors image:Y-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Y-TableImage.png License: GNU Free Documentation License Contributors: Breeze, Moink, Paddy, Paginazero, Saperaud, 2 anonymous edits Image: Y-39-yttrium-sample-crop.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Y-39-yttrium-sample-crop.jpg License: unknown Contributors: User:Mav Image:Mira 1997.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Mira_1997.jpg License: unknown Contributors: Margarita Karovska (Harvard-Smithsonian Center for Astrophysics) and NASA Image:Johan Gadolin.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Johan_Gadolin.jpg License: unknown Contributors: Mav, Nysalor, Tsemii, 竹麦魚(Searobin) Image:Xenotímio1.jpeg Source: http://en.wikipedia.org/w/index.php?title=File:Xenotímio1.jpeg License: Creative Commons Attribution-Sharealike 2.0 Contributors: Zimbres Image:Yttrium 1.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Yttrium_1.jpg License: GNU Free Documentation License Contributors: BLueFiSH.as, Maksim Image:Aperture Grille.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Aperture_Grille.jpg License: Public Domain Contributors: Original uploader was Loongyh at en.wikipedia (Original text : Loongyh (talk)) Image:YBCO-modified.jpg Source: http://en.wikipedia.org/w/index.php?title=File:YBCO-modified.jpg License: unknown Contributors: Akinom, Benjah-bmm27, Puppy8800, Xxxx00
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
18
Zirconium
1
Zirconium 40
yttrium ← zirconium → niobium
Ti ↑
Zr ↓
Hf Periodic Table - Extended Periodic Table
General Name, Symbol, Number
zirconium, Zr, 40
Element category
transition metals
Group, Period, Block
4, 5, d
Appearance
silvery white
Standard atomic weight
91.224(2) g·mol
Electron configuration
[Kr] 4d 5s
−1
2
Electrons per shell
2
2, 8, 18, 10, 2 Physical properties
Phase
solid
Density (near r.t.)
6.52 g·cm
Liquid density at m.p.
5.8 g·cm−3
Melting point
2128 K (1855 °C, 3371 °F)
Boiling point
4682 K (4409 °C, 7968 °F)
Heat of fusion
14 kJ·mol−1
Heat of vaporization
573 kJ·mol−1
Specific heat capacity
(25 °C) 25.36 J·mol−1·K−1
−3
Vapor pressure P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
2639
2891
3197
3575
4053
4678
Atomic properties Crystal structure
hexagonal close-packed
Zirconium
Oxidation states
Electronegativity Ionization energies (more)
2 [1]
4, 3, 2, 1, (amphoteric oxide) 1.33 (Pauling scale) 1st: 640.1 kJ·mol−1 2nd: 1270 kJ·mol−1 3rd: 2218 kJ·mol−1
Atomic radius
160 pm
Covalent radius
175±7 pm Miscellaneous [2]
Magnetic ordering
paramagnetic
Electrical resistivity
(20 °C) 421 n Ω·m
Thermal conductivity
(300 K) 22.6 W·m
Thermal expansion
(25 °C) 5.7 µm·m
Speed of sound (thin rod)
(20 °C) 3800 m/s
Young's modulus
68 GPa
Shear modulus
33 GPa
−1
−1
Poisson ratio
0.34
Mohs hardness
5.0
Vickers hardness
903 MPa
Brinell hardness
650 MPa
CAS registry number
7440-67-7 Most-stable isotopes
−1
·K
·K
−1
Zirconium
3
Main article: Isotopes of zirconium iso 88
Zr
NA syn
half-life 83.4 d
DM ε
-
γ 89
Zr
syn
78.4 h
51.45%
90
91
11.22%
91
92
17.15%
92
Zr Zr Zr
93
Zr
syn
94
17.38%
96
2.8%
Zr Zr
0.392D
DP 88
Y
-
ε
-
89
β+
0.902
89
0.909D
-
−
0.060
93
− −
-
94
− −
3.348
96
γ 90
DE (MeV)
Y Y
Zr is stable with 50 neutron Zr is stable with 51 neutron Zr is stable with 52 neutron 6
1.53×10 y 17
1.1 × 10
β y
19 [3]
2.0×10 y
β β β β
Nb Mo Mo
References
Zirconium (pronounced /zərˈkoʊniəm/) is a chemical element with the symbol Zr and atomic number 40. It is a lustrous, gray-white, strong transition metal that resembles titanium. Zirconium is used as an alloying agent due to its high resistance to corrosion. It is never found as a native metal; it is obtained mainly from the mineral zircon, which can be purified by chlorine. Zirconium was first isolated in an impure form in 1824 by Jöns Jakob Berzelius. Zirconium rod
Zirconium has no known biological role. Zirconium forms both inorganic and organometallic compounds such as zirconium dioxide and zirconocene dichloride, respectively. There are five naturally-occurring isotopes, three of which are stable. Short-term exposure to zirconium powder causes minor irritation, and inhalation of zirconium compounds can cause skin and lung granulomas.
Characteristics Zirconium is a lustrous, grayish-white, soft, ductile, and malleable metal which is solid at room temperature, though it becomes hard and brittle at lower purities.[4] [5] In powder form, zirconium is highly flammable, but the solid form is far less prone to igniting. Zirconium is highly resistant to corrosion by alkalis, acids, salt water, and other agents.[6] However, it will dissolve in hydrochloric and sulfuric acid, especially when fluorine is present.[7] Alloys with zinc become magnetic below 35 K.[6]
Zirconium The melting point of zirconium is at 1855°C, and the boiling point is at 4409°C.[6] Zirconium has an electronegativity of 1.33 on the Pauling scale. Of the elements within d-block, zirconium has the fourth lowest electronegativity after yttrium, lutetium, and hafnium.[8]
Applications Because of zirconium's excellent resistance to corrosion, it is often used as an alloying agent in materials that are exposed to corrosive agents, such as surgical appliances, explosive primers, vacuum tube getters and filaments. Zirconium dioxide (ZrO2) is used in laboratory crucibles, metallurgical furnaces, and as a refractory material.[6] Zircon (ZrSiO4) is cut into gemstones for use in jewelry. Zirconium carbonate (3ZrO2·CO2·H2O) was used in lotions to treat poison ivy, but this was discontinued as it caused bad skin reactions in some cases.[4] 90% of all zirconium produced is used in nuclear reactors because of its low neutron-capture cross-section and resistance to corrosion.[5] [6] Zirconium alloys are used in space vehicle parts for their resistance to heat, an important quality given the extreme heat associated with atmospheric reentry.[9] Zirconium is also a component in some abrasives, such as grinding wheels and sandpaper.[10] Zirconium is used in weapons such as the BLU-97/B Combined Effects Bomb for incendiary effect. Zirconium in the oxidized form is also used in dentistry for crowning of the teeth because of its biocompatibility, strength and appearance.
Refining Upon being collected from coastal waters, the solid mineral zircon is purified by spiral concentrators to remove excess sand and gravel and by magnetic separators to remove ilmenite and rutile. The byproducts can then be dumped back into the water safely, as they are all natural components of beach sand. The refined zircon is then purified into pure zirconium by chlorine or other agents, then sintered until sufficiently ductile for metalworking.[5] Zirconium and hafnium are both contained in zircon and they are quite difficult to separate due to their similar chemical properties.[9]
History The zirconium-containing mineral zircon, or its variations (jargoon, hyacinth, jacinth, ligure), were mentioned in biblical writings.[6] [9] The mineral was not known to contain a new element until 1789,[10] when Klaproth analyzed a jargoon from the island of Sri Zirconium crystal bar, 99,97%, made Lanka in the Indian Ocean. He named the new element by the crystal bar process Zirkonerde (zirconia).[6] Humphry Davy attempted to isolate this new element in 1808 through electrolysis, but failed.[4] Zirconium (from Syriac zargono,[11] Arabic zarkûn from Persian zargûn نوگرز meaning "gold like")[9] was first isolated in an impure form in 1824 by Berzelius by heating a mixture of potassium and potassium-zirconium fluoride in a small decomposition process conducted in an iron tube.[6] The crystal bar process (or Iodide process), discovered by Anton Eduard van Arkel and Jan Hendrik de Boer in 1925, was the first industrial process for the commercial production of pure metallic zirconium. The process involved thermally decomposing zirconium
4
Zirconium
5
tetraiodide. It was superseded in 1945 by the much cheaper Kroll process developed by William Justin Kroll, in which zirconium tetrachloride is broken down by magnesium.[5] [12]
Occurrence Geological
Zirconium output in 2005
Zirconium has a concentration of about 130 mg/kg within the earth's crust and about .026 μg/L in sea water,[13] though it is never found in nature as a native metal. The principal commercial source of zirconium is the zirconium silicate mineral, zircon (ZrSiO4),[4] which is found primarily in Australia, Brazil, India, Russia, South Africa, and the United States, as well as in smaller deposits around the world.[5] 80% of zircon mining occurs in Australia and South Africa.[4] Zircon resources exceed 60 million metric tons worldwide[14] and annual worldwide zirconium production is approximately 900,000 metric tons.[13]
Zircon is a by-product of the mining and processing of the titanium minerals ilmenite and rutile, as well as tin mining.[15] From 2003 to 2007, zircon prices have steadily increased from $360 to $840 per metric ton.[14] World production trend of zirconium mineral concentrates Zirconium also occurs in more than 140 other recognized mineral species including baddeleyite and kosnarite.[16] This metal is commercially produced mostly by the reduction of the zirconium(IV) chloride with magnesium metal in the Kroll process.[6] Commercial-quality zirconium for most uses still has a content of 1% to 3% hafnium.[4] This element is relatively-abundant in S-type stars, and it has been detected in the sun and in meteorites. Lunar rock samples brought back from several Apollo program missions to the moon have a quite high zirconium oxide content relative to terrestrial rocks.[6] See also zirconium minerals.
Biological Zirconium has no known biological role, though zirconium salts are of low toxicity. The human body contains, on average, only 1 milligram of zirconium, and daily intake is approximately 50 μg per day. Zirconium content in human blood is as low as 10 parts per billion. Aquatic plants readily take up soluble zirconium, but it is rare in land plants. 70% of plants have no zirconium content at all, and those that do have as little as 5 parts per billion.[4]
Zirconium
Compounds As a transition metal, zirconium forms various inorganic compounds, such as zirconium dioxide (ZrO2). This compound, also referred to as zirconia, has exceptional fracture toughness and chemical resistance, especially in its cubic form.[17] These properties make zirconia useful as a thermal barrier coating,[18] though it is also a common diamond substitute.[17] Zirconium tungstate is an unusual substance in that it shrinks in all directions when heated, whereas other elements expand when heated.[6] ZrZn2 is one of only two substances to exhibit superconductivity and ferromagnetism simultaneously, with the other being UGe2.[19] Other inorganic zirconium compounds include zirconium(II) hydride, zirconium nitride, and zirconium tetrachloride (ZrCl4), which is used in the Friedel-Crafts reaction.[20] Organozirconium chemistry is the study of compounds containing a carbon-zirconium bond. These organozirconium compounds are often employed as polymerization catalysts. The first such compound was zirconocene dibromide, prepared in 1952 by John M. Birmingham at Harvard University.[21] Schwartz's reagent, prepared in 1970 by P. C. Wailes and H. Weigold,[22] is a metallocene used in organic synthesis for transformations of alkenes and alkynes.[23]
Isotopes Naturally-occurring zirconium is composed of five isotopes. 90Zr, 91Zr, and 92Zr are stable. 94 Zr has a half-life of 1.10×1017 years. 96Zr has a half-life of 2.4×1019 years, making it the longest-lived radioisotope of zirconium. Of these natural isotopes, 90Zr is the most common, making up 51.45% of all zirconium. 96Zr is the least common, comprising only 2.80% of zirconium.[24] 28 artificial isotopes of zirconium have been synthesized, ranging in atomic mass from 78 to 110. 93Zr is the longest-lived artificial isotope, with a half-life of 1.53×106 years. 110Zr, the heaviest isotope of zirconium, is also the shortest-lived, with an estimated half-life of only 30 milliseconds. Radioactive isotopes at or above mass number 93 decay by β−, whereas those at or below 89 decay by β+. The only exception is 88Zr, which decays by ε.[24] Zirconium also has six metastable isomers, 83mZr, 85mZr, 89mZr, 90m1Zr, 90m2Zr, and 91mZr. Of these, 90m2Zr has the shortest half-life at 131 nanoseconds. 89mZr is the longest lived with a half-life of 4.161 minutes.[24]
Toxicity Short-term exposure to zirconium powder can cause irritation, but only contact with the eyes requires medical attention.[25] Inhalation of zirconium compounds can cause skin and lung granulomas. Zirconium aerosols can cause pulmonary granulomas. Persistent exposure to zirconium tetrachloride resulted in increased mortality in rats and guinea pigs and a decrease of blood hemoglobin and red blood cells in dogs. OSHA recommends a 5 mg/m3 time weighted average limit and a 10 mg/m3 short-term exposure limit.[26]
6
Zirconium
7
See also • Zirconium compounds • Zirconium minerals
External links • WebElements.com: Zirconium
[27]
pnb:مینوکرز
References [1] " Zirconium: zirconium(I) fluoride compound data (http:/ / openmopac. net/ data_normal/ zirconium(i) fluoride_jmol. html)". OpenMOPAC.net. . Retrieved 2007-12-10. [2] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [3] Pritychenko, Boris; V. Tretyak. " Adopted Double Beta Decay Data (http:/ / www. nndc. bnl. gov/ bbdecay/ list. html)". National Nuclear Data Center. . Retrieved 2008-02-11. [4] Emsley, John (2001). Nature's Building Blocks. Oxford: Oxford University Press. pp. 506–510. ISBN 0-19-850341-5. [5] " Zirconium (http:/ / www. madehow. com/ Volume-1/ Zirconium. html)". How Products Are Made. Advameg Inc.. 2007. . Retrieved 2008-03-26. [6] Lide, David R., ed. (2007–2008), "Zirconium", CRC Handbook of Chemistry and Physics, 4, New York: CRC Press, pp. 42, 978-0-8493-0488-0 [7] Considine, Glenn D., ed. (2005), "Zirconium", Van Nostrand's Encyclopedia of Chemistry, New York: Wylie-Interscience, pp. 1778–1779, ISBN 0-471-61525-0 [8] Winter, Mark (2007). " Electronegativity (Pauling) (http:/ / www. webelements. com/ webelements/ properties/ text/ image-flash/ electroneg-pauling. html)". University of Sheffield. . Retrieved 2008-03-05. [9] Stwertka, Albert (1996). A Guide to the Elements. Oxford University Press. pp. 117–119. ISBN 0-19-508083-1. [10] Krebs, Robert E. (1998). The History and Use of our Earth's Chemical Elements. Westport, Connecticut: Greenwood Press. pp. 98–100. ISBN 0-313-30123-9. [11] Pearse, Roger (2002-09-16). " Syriac Literature (http:/ / www. tertullian. org/ rpearse/ oriental/ syriac. htm)". . Retrieved 2008-02-11. [12] Hedrick, James B. (1998), " Zirconium (http:/ / minerals. usgs. gov/ minerals/ pubs/ metal_prices/ metal_prices1998. pdf)" (PDF), Metal Prices in the United States through 1998, US Geological Survey, pp. 175–178, , retrieved 2008-02-26 [13] Peterson, John; MacDonell, Margaret (2007), " Zirconium (http:/ / www. evs. anl. gov/ pub/ doc/ ANL_ContaminantFactSheets_All_070418. pdf)" (PDF), Radiological and Chemical Fact Sheets to Support Health Risk Analyses for Contaminated Areas, Argonne National Laboratory, pp. 64–65, , retrieved 2008-02-26 [14] " Zirconium and Hafnium (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ zirconium/ mcs-2008-zirco. pdf)" (PDF). Mineral Commodity Summaries (US Geological Survey): 192–193. January 2008. . Retrieved 2008-02-24. [15] Callaghan, R. (2008-02-21). " Zirconium and Hafnium Statistics and Information (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ zirconium/ )". US Geological Survey. . Retrieved 2008-02-24. [16] Ralph, Jolyon; Ida Ralph (2008). " Minerals that include Zr (http:/ / www. mindat. org/ chemsearch. php?inc=Zr,& exc=& sub=Search+ for+ Minerals)". Mindat.org. . Retrieved 2008-02-23. [17] " Zirconia (http:/ / www. azom. com/ details. asp?ArticleID=133#_Key_Properties)". AZoM.com. 2008. . Retrieved 2008-03-17. [18] Gauthier, V.; Dettenwanger, F.; Schütze, M. (2002-04-10). "Oxidation behavior of γ-TiAl coated with zirconia thermal barriers". Intermetallics (Frankfurt, Germany: Karl Winnacker Institut der Dechema) 10 (7): 667–674. doi: 10.1016/S0966-9795(02)00036-5 (http:/ / dx. doi. org/ 10. 1016/ S0966-9795(02)00036-5). [19] Day, Charles (September 2001). "Second Material Found that Superconducts in a Ferromagnetic State". Physics Today (American Institute of Physics) 54 (9): 16. doi: 10.1063/1.1420499 (http:/ / dx. doi. org/ 10. 1063/ 1. 1420499). [20] Bora U. (2003). "Zirconium Tetrachloride". Synlett: 1073–1074. doi: 10.1055/s-2003-39323 (http:/ / dx. doi. org/ 10. 1055/ s-2003-39323). [21] Rouhi, A. Maureen (2004-04-19). " Organozirconium Chemistry Arrives (http:/ / pubs. acs. org/ cen/ nlw/ 8216sci1. html)". Science & Technology (Chemical & Engineering News) 82 (16): 36–39. ISSN 0009-2347
Zirconium (http:/ / worldcat. org/ issn/ 0009-2347). . Retrieved 2008-03-17. [22] P. C. Wailes and H. Weigold (1970). "Hydrido complexes of zirconium I. Preparation". Journal of Organometallic Chemistry 24: 405–411. doi: 10.1016/S0022-328X(00)80281-8 (http:/ / dx. doi. org/ 10. 1016/ S0022-328X(00)80281-8). [23] D. W. Hart and J. Schwartz (1974). "Hydrozirconation. Organic Synthesis via Organozirconium Intermediates. Synthesis and Rearrangement of Alkylzirconium(IV) Complexes and Their Reaction with Electrophiles". J. Am. Chem. Soc. 96 (26): 8115–8116. doi: 10.1021/ja00833a048 (http:/ / dx. doi. org/ 10. 1021/ ja00833a048). [24] Audi, G (2003). "Nubase2003 Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3–128. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [25] " Zirconium (http:/ / www. oit. org/ public/ english/ protection/ safework/ cis/ products/ icsc/ dtasht/ _icsc14/ icsc1405. htm)", International Chemical Safety Cards, International Labour Organization, October 2004, , retrieved 2008-03-30 [26] " Zirconium Compounds (http:/ / www. cdc. gov/ niosh/ pel88/ 7440-67. html)". National Institute for Occupational Health and Safety. 2007-12-17. . Retrieved 2008-02-17. [27] http:/ / webelements. com/ zirconium/
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Article Sources and Contributors
Article Sources and Contributors Zirconium Source: http://en.wikipedia.org/w/index.php?oldid=308579942 Contributors: 777B, Ahoerstemeier, Alansohn, Alchemist-hp, Algebraist, AlimanRuna, Anclation, Andres, Antandrus, Anwar saadat, Archimerged, ArgentTurquoise, Atarr, AxelBoldt, Baccyak4H, Beetstra, Bencherlite, Bender235, BendersGame, Bgt5bhu8, BlueEarth, Bryan Derksen, CYD, Cacahueten, Callipides, Calvin 1998, Can't sleep, clown will eat me, CanadianLinuxUser, CanisRufus, Carnildo, Catgut, Chaosfeary, Chem-awb, Cjnm, Conversion script, Coppertwig, Cristina0530, Crusader75, Cryptic C62, Crüsäder, Ctesiphon7, Dajwilkinson, Dana boomer, Darrien, David Latapie, Delldot, Delta G, DerHexer, Donarreiskoffer, Download, Dschwen, DuncanHill, Dwmyers, EPO, Edgar181, El C, EliasAlucard, Emperorbma, Epbr123, EricV89, Evertype, Everyking, FelisLeo, Femto, Firq, Fnielsen, Franamax, FreplySpang, Giftlite, Grendelkhan, Guaca, Hadal, Hak-kâ-ngìn, Haligonian1, Hannibal, Hydrogen Iodide, I-hunter, Ianml, Icairns, Ideyal, ImperatorExercitus, It Is Me Here, Ixfd64, J.delanoy, JWB, Jacob.l345, Jaraalbe, JdH, Joanjoc, John, Jose77, Juliancolton, Kakofonous, Karl-Henner, Karlhahn, KathrynLybarger, Kbh3rd, Ketsuekigata, Kristen Eriksen, Ktsquare, Kurykh, Kwamikagami, Kyle Barbour, LA2, LeoNomis, Leyo, Lst27, LuigiManiac, Luna Santin, MER-C, Martial75, Materialscientist, Mav, Megan1967, Mejor Los Indios, Mgimpel, Michael Devore, Midgley, Mike5904, Minesweeper, Mjager, Mor, Mouser, Mycroft7, Narayanese, Narge, NawlinWiki, Nergaal, NewEnglandYankee, Nibuod, Night Gyr, Nihiltres, Nn123645, Open2universe, Physchim62, PierreAbbat, Pinkadelica, Plasmic Physics, PlatinumX, Plexust, Poolkris, RTC, Remember, RexNL, Riana, Rifleman 82, Rjwilmsi, Rmrfstar, Roberta F., Romanm, Rtcoles, Saperaud, Schneelocke, Scohoust, Scoutersig, Sea Dagon, Sengkang, Sfuerst, Shaddack, Shiraun, SidP, Sjakkalle, Sl, Smallweed, Smarterdude678, Smartperson678, Soliloquial, SpK, Sprintstar, Stephenb, Stifynsemons, Stone, Sunborn, Tagishsimon, Tetracube, Tfl, Thingg, Thricecube, Tiger Khan, Timeineurope, V1adis1av, Versus22, Vsmith, Walkerma, Walton One, Warut, Wetman, Where, Willking1979, Xagent86, Xenowiki, Yekrats, Yyy, Zeldafanzunite, Ψ-113μ, 394 anonymous edits
Image Sources, Licenses and Contributors image:Zr-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Zr-TableImage.png License: GNU Free Documentation License Contributors: Joanjoc, Paddy, Saperaud, 1 anonymous edits Image: Zr,40.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Zr,40.jpg License: GNU Free Documentation License Contributors: Joanjoc, Paginazero, Saperaud Image:Zirconium rod.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Zirconium_rod.jpg License: unknown Contributors: User:Dschwen File:Zr-crystal-bar.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Zr-crystal-bar.jpg License: Creative Commons Attribution-Sharealike 3.0 Germany Contributors: User:Alchemist-hp Image:2005zirconium.PNG Source: http://en.wikipedia.org/w/index.php?title=File:2005zirconium.PNG License: Public Domain Contributors: User:Anwar_saadat/bubble_maps_(FAQ) Image:Zirconium mineral concentrates - world production trend.svg Source: http://en.wikipedia.org/w/index.php?title=File:Zirconium_mineral_concentrates_-_world_production_trend.svg License: Public Domain Contributors: User:Leyo
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
9
Niobium
1
Niobium 41
zirconium ← niobium → molybdenum
V ↑
Nb ↓
Ta Periodic Table - Extended Periodic Table
General Name, Symbol, Number
niobium, Nb, 41
Element category
transition metals
Group, Period, Block
5, 5, d
Appearance
gray metallic
Standard atomic weight
92.90638(2) g·mol
Electron configuration
[Kr] 4d 5s
−1
4
Electrons per shell
1
2, 8, 18, 12, 1 Physical properties
Phase
solid
Density (near r.t.)
8.57 g·cm−3
Melting point
2750 K (2477 °C, 4491 °F)
Boiling point
5017 K (4744 °C, 8571 °F)
Heat of fusion
30 kJ·mol−1
Heat of vaporization
689.9 kJ·mol−1
Specific heat capacity
(25 °C) 24.60 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
2942
3207
3524
3910
4393
5013
Atomic properties Crystal structure
body centered cubic
Niobium
2
Oxidation states
5, 4, 3, 2, -1 (mildly acidic oxide)
Electronegativity
1.6 (Pauling scale)
Ionization energies (more)
1st: 652.1 kJ·mol−1 2nd: 1380 kJ·mol−1 3rd: 2416 kJ·mol−1
Atomic radius
146 pm
Covalent radius
164±6 pm Miscellaneous
Magnetic ordering
paramagnetic
Electrical resistivity
(0 °C) 152 nΩ·m
Thermal conductivity
(300 K) 53.7 W·m
Thermal expansion
(25 °C) 7.3 µm·m
Speed of sound (thin rod)
(20 °C) 3480 m/s
Young's modulus
105 GPa
Shear modulus
38 GPa
Bulk modulus
170 GPa
−1
−1
Poisson ratio
0.40
Mohs hardness
6.0
Vickers hardness
1320 MPa
Brinell hardness
736 MPa
CAS registry number
7440-03-1 Most-stable isotopes
−1
·K
·K
−1
Niobium
3
Main article: Isotopes of niobium iso
NA
half-life
DM
DE (MeV)
syn
6.8×102 y
ε
91m
syn
60.86 d
IT
92
syn
10.15 d
ε
-
γ
0.934
ε
-
γ
0.561, 0.934
91
Nb Nb
Nb
92
Nb
93
Nb
syn
100%
93m
syn
94
syn
Nb
Nb
95
Nb
95m
Nb
syn
syn
3.47×107y
0.104e
DP 91
Zr
91
Nb
92
Zr
92
Zr
-
93
Nb is stable with 52 neutron
16.13 y
IT 4
2.03×10 y
34.991 d
3.61 d
0.031e
93
β
0.471
94
γ
0.702, 0.871
−
Nb Mo
-
β
0.159
95
γ
0.765
-
IT
0.235
95
−
Mo
Nb
References
Niobium (pronounced /naɪˈoʊbiəm/) (Greek mythology: Niobe, daughter of Tantalus), or columbium (/kəˈlʌmbiəm/), is the chemical element with the symbol Nb and the atomic number 41. A rare, soft, grey, ductile transition metal, niobium is found in the minerals pyrochlore, the main commercial source for niobium, and columbite. Niobium has physical and chemical properties similar to those of the element tantalum, and the two are therefore difficult to distinguish. The English chemist Charles Hatchett reported a new element similar to tantalum in 1801, and named it columbium. In 1809, the English chemist William Hyde Wollaston wrongly concluded that tantalum and columbium were identical. The German chemist Heinrich Rose determined in 1846 that tantalum ores contain a second element, which he named niobium. In 1864 and 1865, a series of scientific findings clarified that niobium and columbium were the same element (as distinguished from tantalum), and for a century both names were used interchangeably. The name of the element was officially adopted as niobium in 1949. It was not until the early 20th century that niobium was first used commercially. Brazil is the leading producer of niobium and ferroniobium, an alloy of niobium and iron. Niobium is used mostly in alloys, the largest part in special steel such as that used in gas pipelines. Although alloys contain only a maximum of 0.1%, that small percentage of Niobium improves the strength of the steel. The temperature stability of niobium-containing superalloys is important for its use in jet engines and rocket engines. Niobium is used in various superconducting materials. These superconducting alloys, also containing titanium and tin, are widely used in the superconducting magnets of MRI scanners. Other applications of niobium include its use in welding, nuclear industries, electronics, optics, numismatics and jewellery. In the last two applications, niobium's low toxicity and ability to
Niobium
4
be coloured by anodisation are particular advantages.
History Niobium was discovered by the English chemist Charles Hatchett in 1801.[1] He found a new element in a mineral sample that had been sent to England from Massachusetts, United States in 1734 by a John Winthrop,[2] and named the mineral columbite and the new element columbium after Columbia, the poetical name for America.[3] The columbium discovered by Hatchett was probably a mixture of the new element with tantalum.[3]
Charles Hatchett discoverer of columbium.
Picture of a Hellenistic sculpture representing Niobe by Giorgio Sommer
Subsequently, there was considerable confusion[4] over the difference between columbium (niobium) and the closely related tantalum. In 1809, the English chemist William Hyde Wollaston compared the oxides derived from both columbium—columbite, with a density 5.918 g/cm3, and tantalum—tantalite, with a density 7.935 g/cm3, and concluded that the two oxides, despite the significant difference in density, were identical; thus he kept the name tantalum.[4] This conclusion was disputed in 1846 by the German chemist Heinrich Rose, who argued that there were two different elements in the tantalite sample, and named them after children of Tantalus: niobium (from Niobe), and pelopium (from Pelops).[5] [6] This confusion arose from the minimal observed differences between tantalum and niobium. Both tantalum and niobium react with chlorine and traces of oxygen, including atmospheric concentrations, with niobium forming two compounds: the white volatile niobium pentachloride (NbCl5) and the non-volatile niobium oxychloride (NbOCl3). The claimed new elements pelopium, ilmenium and dianium[7] were in fact identical to niobium or mixtures of niobium and tantalum.[8]
The differences between tantalum and niobium were unequivocally demonstrated in 1864 by Christian Wilhelm Blomstrand,[8] and Henri Etienne Sainte-Claire Deville, as well as Louis J. Troost, who determined the formulas of some of the compounds in 1865[8] [9] and finally by the Swiss chemist Jean Charles Galissard de Marignac[10] in 1866, who all proved that there were only two elements. These discoveries did not stop scientists from publishing articles about ilmenium until 1871.[11] De Marignac was the first to prepare the metal in 1864, when he reduced niobium chloride by heating it in an atmosphere of hydrogen.[12] Although de Marignac was able to produce tantalum-free niobium on a larger scale by 1866, it was not until the early 20th century that niobium was first used commercially, in incandescent lamp filaments.[9] This use quickly became obsolete through the replacement of niobium with tungsten, which has a higher melting point and thus is preferable for use in incandescent lamps. The discovery that niobium improves the strength of steel was made in the 1920s, and this remains its predominant use.[9] In 1961 the American physicist Eugene Kunzler and coworkers at Bell Labs discovered that niobium-tin continues to exhibit superconductivity in the presence of strong electric currents and magnetic fields,[13]
Niobium
5
making it the first material known to support the high currents and fields necessary for making useful high-power magnets and electrically powered machinery. This discovery would allow—two decades later—the production of long multi-strand cables that could be wound into coils to create large, powerful electromagnets for rotating machinery, particle [14] [15] accelerators, or particle detectors. Columbium (symbol Cb[16] ) was the name originally given to this element by Hatchett, and this name remained in use in American journals—the last paper published by American Chemical Society with columbium in its title dates from 1953[17] —while niobium was used in Europe. To end this confusion, the name niobium was chosen for element 41 at the 15th Conference of the Union of Chemistry in Amsterdam in 1949.[18] A year later this name was officially adopted by the International Union of Pure and Applied Chemistry (IUPAC) after 100 years of controversy, despite the chronological precedence of the name Columbium.[18] The latter name is still sometimes used in US industry.[19] This was a compromise of sorts;[18] the IUPAC accepted tungsten instead of wolfram, in deference to North American usage; and niobium instead of columbium, in deference to European usage. Not everyone agreed, and while many leading chemical societies and government organizations refer to it by the official IUPAC name, many leading metallurgists, metal societies, and the United States Geological Survey still refer to the metal by the original "columbium".[20] [21]
Characteristics Niobium is a lustrous, grey, ductile, paramagnetic metal in group 5 of the periodic table (see table to right), Z
Element
No. of electrons/shell
23
vanadium
2, 8, 11, 2
41
niobium
2, 8, 18, 12, 1
73
tantalum
2, 8, 18, 32, 11, 2
105
dubnium
2, 8, 18, 32, 32, 11, 2
although it has an atypical configuration in its outermost electron shells compared to the rest of the members. (This can be observed in the neighborhood of niobium (41), ruthenium (44), rhodium (45), and palladium (46).) The metal takes on a bluish tinge when exposed to air at room temperature for extended periods.[22] Despite presenting a high melting point in elemental form (2,468 °C), it has a low density in comparison to other refractory metals. Furthermore, it is corrosion resistant, exhibits superconductivity properties, and forms dielectric oxide layers. These properties— especially the superconductivity —are strongly dependent on the purity of the niobium metal.[23] When very pure, it is comparatively soft and ductile, but impurities make it harder.[24] The atoms of niobium is slightly less electropositive and smaller than the atoms of its predecessor in the periodic table, zirconium, while it is virtually identical in size to the heavier tantalum atoms as a consequence of the lanthanide contraction.[24] As a result, niobium's chemical properties are very similar to the chemical properties of tantalum, which appears directly below niobium in the periodic table.[9] Although its corrosion resistance is not as outstanding as that of tantalum, its lower price and greater availability make niobium attractive for less exact uses such as linings in chemical plants.[24]
Niobium
Isotopes Naturally occurring niobium is composed of one stable isotope, 93Nb.[25] As of 2003, at least 32 radioisotopes have also been synthesized, ranging in atomic mass from 81 to 113. The most stable of these is 92Nb with a half-life of 34.7 million years. One of the least stable is 113Nb, with an estimated half-life of 30 milliseconds. Isotopes that are lighter than the stable 93Nb tend to decay by β+ decay, and those that are heavier tend to decay by β- decay, with some exceptions. 81Nb, 82Nb, and 84Nb have minor β+ delayed proton emission decay paths, 91Nb decays by electron capture and positron emission, and 92Nb decays by both β+ and β- decay.[25] At least 25 nuclear isomers have been described, ranging in atomic mass from 84 to 104. Within this range, only 96Nb, 101Nb, and 103Nb do not have isomers. The most stable of niobium's isomers is 93mNb with a half-life of 16.13 years. The least stable isomer is 84mNb with a half-life of 103 ns. All of niobium's isomers decay by isomeric transition or beta decay except 92m1Nb, which has a minor electron capture decay chain.[25]
Chemistry Niobium is in many ways similar to its predecessors in group 5. It reacts with most nonmetals at high temperatures: niobium reacts with fluorine at room temperature, with chlorine and hydrogen at 200 °C, and with nitrogen at 400 °C, giving products that are frequently interstitial and nonstoichiometric.[24] The metal begins to oxidize in air at 200 °C,[26] and is resistant to corrosion by fused alkalis and by acids, including aqua regia, hydrochloric, sulfuric, nitric and phosphoric acids.[24] Niobium is attacked by hot, concentrated mineral acids, such as fluorhydric acid and fluorhydric/nitric acid mixtures. Although niobium exhibits all the formal oxidation states from +5 down to -1, its most stable state is +5.[24] Niobium is able to form oxides with the oxidation states +5 (Nb2O5), +4 (NbO2) and +3 (Nb2O3),[26] as well as with the rarer oxidation state +2 (NbO).[27] The most stable oxidation state is +5, the pentoxide which, along with the dark green non-stoichiometric dioxide, is the most common of the oxides.[26] Niobium pentoxide is used mainly in the production of capacitors, optical glass, and as starting material for several niobium compounds.[28] The compounds are created by dissolving the pentoxide in basic hydroxide solutions or by melting it in another metal oxide. Examples are lithium niobate (LiNbO3) and lanthanum niobate (LnNbO4). In the lithium niobate, the niobate ion NbO3− is not alone but part of a trigonally distorted perovskite-like structure, while the lanthanum niobate contains lone NbO43− ions.[26] Lithium niobate, which is a ferroelectric, is used extensively in mobile telephones and optical modulators, and for the manufacture of surface acoustic wave devices. It belongs to the ABO3 structure ferroelectrics like lithium tantalate and barium titanate.[29]
6
Niobium
Niobium forms halogen compounds in the oxidation states of +5, +4, and +3 of the type NbX5, NbX4, and NbX3, although multi-core complexes and [26] [30] substoichiometric compounds are also formed. Niobium pentafluoride (NbF5) is a white solid with a melting point of 79.0 °C and niobium pentachloride (NbCl5) is a yellowish-white solid (see image at left) with a melting point of 203.4 °C. Both are hydrolyzed by water and react with additional niobium at elevated temperatures by forming the black and highly hygroscopic niobium tetrafluoride (NbF4) and niobium tetrachloride (NbCl4). While the trihalogen compounds Niobium pentachloride (NbCl5) can be obtained by reduction of the pentahalogens with hydrogen, the dihalogen compounds do not exist.[26] Spectroscopically, the monochloride (NbCl) has been observed at high temperatures.[31] The fluorides of niobium can be used after its separation from tantalum.[32] The niobium pentachloride is used in organic chemistry as a Lewis acid in activating alkenes for the carbonyl-ene reaction and the Diels-Alder reaction.[33] The pentachloride is also used to generate the organometallic compound niobocene dichloride ((C5H5)2NbCl2), which in turn is used as a starting material for other organoniobium compounds.[34] Other binary compounds of niobium include niobium nitride (NbN), which becomes a superconductor at low temperatures and is used in detectors for infrared light,[35] and niobium carbide, an extremely hard, refractory, ceramic material, commercially used in tool bits for cutting tools. The compounds niobium-germanium (Nb3Ge) and niobium-tin (Nb3Sn), as well as the niobium-titanium alloy, are used as a type II superconductor wire for superconducting magnets.[36] [37] Niobium sulfide as well as a few interstitial compounds of niobium with silicon are also known.[24]
Occurrence According to estimates, niobium is 33rd on the list of the most common elements in the Earth’s crust with 20 ppm.[38] The abundance on Earth should be much greater, but the “missing” niobium may be located in the Earth’s core due to the metal's high density.[20] The free element is not found in nature, but it does occur in minerals.[24] Minerals that contain niobium often also contain tantalum, for example, columbite ((Fe,Mn)(Nb,Ta)2O6), columbite-tantalite (or coltan, (Fe,Mn)(Ta,Nb)2O6) and pyrochlore [32] ((Na,Ca)2Nb2O6(OH,F)). Columbite-tantalite minerals are most usually found as accessory minerals in pegmatite intrusions, and in alkaline intrusive rocks. Less common are the niobates of calcium, uranium, thorium and the rare earth elements such as pyrochlore and euxenite ((Y,Ca,Ce,U,Th)(Nb,Ta,Ti)2O6). These large deposits of niobium have been found associated with carbonatites (carbonate-silicate igneous rocks) and as a constituent of pyrochlore.[39] The two largest deposits of pyrochlore were found in the 1950s in Brazil and Canada, and both countries are still the major producers of niobium mineral concentrates.[9] The largest deposit is hosted within a carbonatite intrusion at Araxá, Minas Gerais Brazil, owned by CBMM (Companhia Brasileira de Metalurgia e Mineração); the other deposit is located at Catalão, Goiás owned by Anglo American plc (through its subsidiary Mineração Catalão),
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also hosted within a carbonatite intrusion.[40] Altogether these two Brazilian mines produce around 75% of world supply. The third largest producer of niobium is the carbonatite-hosted Niobec Mine, Saint-Honoré near Chicoutimi, Quebec owned by Iamgold Corporation Ltd, which produces around 7% of world supply.[40] Extensive though unexploited resources are located in Nigeria, Democratic Republic of Congo, Tanzania, Malawi, Australia and Russia.
Production After the separation from the other minerals, the mixed oxides of tantalum Ta2O5 and niobium Nb2O5 are obtained. The first step in the processing is the reaction of the oxides with hydrofluoric [32] acid: Ta2O5 + 14HF → 2H2[TaF7] + 5H2O, and Nb2O5 + 10HF 2H2[NbOF5] + 3H2O
Niobium producers in 2007
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The first industrial scale separation, developed by de Marignac, used the difference in solubility between the complex niobium and tantalum fluorides, dipotassium oxypentafluoroniobate monohydrate (K2[NbOF5]·H2O) and dipotassium heptafluorotantalate (K2[TaF7]) in water. Newer processes use the liquid extraction of the fluorides from aqueous solution by organic solvents like cyclohexanone.[32] The complex niobium and tantalum fluorides are extracted separately from the organic solvent with water and either precipitated by the addition of potassium fluoride to produce a potassium fluoride complex, or precipitated with ammonia as the pentoxide:[26] H2[NbOF5] + 2KF → K2[NbOF5]↓ + 2HF, then 2H2[NbOF5] + 10NH4OH → Nb2O5↓ + 10NH4F + 7H2O Several methods are used for the reduction to metallic niobium. The electrolysis of a molten mixture of K2[NbOF5] and sodium chloride is one; the other is the reduction of the fluoride with sodium. With this method niobium with a relatively high purity can be obtained. In large scale production the reduction of Nb2O5 with hydrogen or carbon[26] is used. In the process involving the aluminothermic reaction a mixture of iron oxide and niobium oxide is reacted with aluminium: 3Nb2O5 + Fe2O3 + 12Al → 6Nb + 2Fe + 6Al2O3 To enhance the reaction, small amounts of oxidizers like sodium nitrate are added. The result is aluminium oxide and ferroniobium, an alloy of iron and niobium used in the steel production.[41] [42] The ferroniobium contains between 60 and 70% of niobium.[40] Without addition of iron oxide, aluminothermic process is used for the production of niobium. Further purification is necessary to reach the grade for superconductive alloys. Electron beam melting under vacuum is the method used by the two major distributors of niobium.[30] [43]
Niobium The United States Geological Survey estimates that the production increased from 38,700 metric tonnes in 2005 to 44,500 tonnes in 2006.[44] [45] The world wide resources are estimated to be 4,400,000 tonnes.[45] During the ten year period between 1995 and 2005, the production more than doubled starting from 17,800 tonnes in 1995.[46]
Applications It is estimated that out of 44,500 metric tons of niobium mined in 2006, 90% ended up in the production of high-grade structural steel, followed by its use in superalloys.[47] The use of niobium alloys for superconductors and in electronic components account only for a small share of the production.[47]
Steel production Niobium is an effective microalloying element for steel. Adding niobium to the steel causes the formation of A niobium foil niobium carbide and niobium nitride within the structure of the steel.[20] These compounds improve the grain refining, retardation of recrystallization, and precipitation hardening of the steel. These effects in turn increase the toughness, strength, formability, and weldability of the microalloyed steel.[20] Microalloyed stainless steels have a niobium content of less than 0.1%.[48] It is an important alloy addition to high strength low alloy steels which are widely used as structural components in modern automobiles.[20] These niobium containing alloys are strong and are often used in pipeline construction.[49] [50]
Superalloys Appreciable amounts of the element, either in its pure form or in the form of high-purity ferroniobium and nickel niobium, are used in nickel-, cobalt-, and iron-base superalloys for such applications as jet engine components, gas turbines, rocket subassemblies, and heat resisting and combustion equipment. Niobium precipitates a hardening γ''-phase within the grain structure of the superalloy.[51] The alloys contain up to 6.5% niobium.[48] One example of a nickel-based niobium-containing superalloy is Inconel 718, which Apollo CSM with the dark rocket nozzle made from niobium-titanium consists of roughly 50% nickel, 18.6% chromium, 18.5% alloy iron, 5% niobium, 3.1% molybdenum, 0.9% titanium, [52] [53] and 0.4% aluminium. These superalloys are used, for example, in advanced air frame systems such as those used in the Gemini program. An alloy used for liquid rocket thruster nozzles, such as in the main engine of the Apollo Lunar Modules, is C103, which consists of 89% niobium, 10% hafnium and 1% titanium.[54] Another niobium alloy was used for the nozzle of the Apollo Service Module. As niobium is oxidized at temperatures above 400 °C, a protective coating is necessary for these
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applications to prevent the alloy from becoming brittle.[54]
Superconducting magnets Niobium becomes a superconductor when lowered to cryogenic temperatures. At atmospheric pressure, it has the highest critical temperature of the elemental superconductors: 9.2 K.[55] Niobium has the largest magnetic penetration depth of any element.[55] In addition, it is one of the three elemental Type II superconductors, along with vanadium and technetium. Niobium-tin and niobium-titanium alloys are used as wires for superconducting magnets capable of A 3 tesla clinical magnetic resonance producing exceedingly strong magnetic fields. These imaging scanner using superconducting magnets are used in magnetic niobium-superconducting alloy resonance imaging and nuclear magnetic resonance instruments as well as in particle accelerators.[56] For example, the Large Hadron Collider uses 600 metric tons of superconducting strands, while the International Thermonuclear Experimental Reactor is estimated to use 600 metric tonnes of Nb3Sn strands and 250 metric tonnes of NbTi strands.[57] In 1992 alone, niobium-titanium wires were used to construct more than 1 billion US dollars worth of clinical magnetic resonance imaging systems.[14]
Numismatics Niobium is used as a precious metal in commemorative coins, often with silver or gold. For example, Austria produced a series of silver niobium euro coins starting in 2003; the colour in these coins is created by diffraction of light by a thin oxide layer produced by anodising.[58] In 2008, six coins are available showing a broad variety of colours in the centre of the coin: blue, green, brown, purple, violet, or yellow. Two more examples are the 2004 Austrian €25 150 Years Semmering Alpine Railway commemorative coin,[59] and the 2006 Austrian €25 European Satellite Navigation commemorative coin.[60] Latvia produced a similar series of coins starting in 2004,[61] with one following in 2007.[62]
A 150 Years Semmering Alpine Railway Coin made of niobium and silver
Other uses Niobium and some niobium alloys are used in medical devices such as pacemakers, because they are physiologically inert (and thus hypoallergenic).[63] Niobium treated with sodium hydroxide forms a porous layer that aids osseointegration.[64] Along with titanium, tantalum, and aluminium, niobium can also be electrically heated and anodized, resulting in a wide array of colours using a process known as reactive metal anodizing which is useful
Niobium in making jewelry.[65] jewelry.[67]
11 [66]
The fact that niobium is hypoallergenic also benefits its use in
The arc-tube seals of high pressure sodium vapor lamps are made from niobium, or niobium with 1% of zirconium, because niobium has a very similar coefficient of thermal expansion to the sintered alumina arc tube ceramic, a translucent material which resists chemical attack or reduction by the hot liquid sodium and sodium vapour contained inside the operating lamp.[68] [69] [70] The metal is also used in arc welding rods for some stabilized grades of stainless steel.[71] Niobium was evaluated as a cheaper alternative to tantalum in capacitors,[72] but tantalum capacitors are still predominant. Niobium is added to glass in order to attain a higher refractive index, a property of use to the optical industry in making thinner corrective glasses. The metal has a low capture cross-section for thermal neutrons;[73] thus it is used in the nuclear industries.[74] The Superconducting Radio Frequency (RF) cavities used in the free electron lasers TESLA and XFEL are made from pure niobium.[75] The high sensitivity of superconducting niobium nitride bolometers make them an ideal detector for electromagnetic radiation in the THz frequency band. These detectors were tested at the Heinrich Hertz Submillimeter Telescope, the South Pole Telescope, the Receiver Lab Telescope, and at APEX and are now used in the HIFI instrument on board the Herschel Space Observatory.[76]
Precautions Niobium has no known biological role. While niobium dust is an eye and skin irritant and a potential fire hazard, elemental niobium on a larger scale is physiologically inert (and thus hypoallergenic) and harmless. It is frequently used in jewelry and has been tested for use in some medical implants.[77] [78] Niobium-containing compounds are rarely encountered by most people, but some are toxic and should be treated with care. The short and long term exposure to niobates and niobium chloride, two chemicals that are water soluble, have been tested in rats. Rats treated with a single injection of niobium pentachloride or niobates show a median lethal dose (LD50) between 10 and 100 mg/kg.[79] [80] [81] For oral administration the toxicity is lower; a study with rats yielded a LD50 after seven days of 940 mg/kg.[79]
External links • • • •
Los Alamos National Laboratory – Niobium [82] WebElements.com – Niobium [83] Tantalum-Niobium International Study Center [84] Niobium for particle accelerators eg ILC. 2005 [85]
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[53] Pottlacher, G.; Hosaeus, H.; Wilthan, B.; Kaschnitz, E.; Seifter, A. (2002). "Thermophysikalische Eigenschaften von festem und flüssigem Inconel 718" (in German). Thermochimica Acta 382: 55–267. doi: 10.1016/S0040-6031(01)00751-1 (http:/ / dx. doi. org/ 10. 1016/ S0040-6031(01)00751-1). [54] Hebda, John (2001-05-02). " Niobium alloys and high Temperature Applications (http:/ / www. cbmm. com. br/ portug/ sources/ techlib/ science_techno/ table_content/ sub_3/ images/ pdfs/ 016. pdf)" (pdf). Niobium Science & Technology: Proceedings of the International Symposium Niobium 2001 (Orlando, Florida, USA) (Companhia Brasileira de Metalurgia e Mineração). . [55] Peiniger, M.; Piel, H. (1985). "A Superconducting Nb3Sn Coated Multicell Accelerating Cavity". Nuclear Science 32 (5): 3610. doi: 10.1109/TNS.1985.4334443 (http:/ / dx. doi. org/ 10. 1109/ TNS. 1985. 4334443). [56] Glowacki, B. A.; Yan, X. -Y.; Fray, D.; Chen, G.; Majoros, M.; Shi, Y. (2002). "Niobium based intermetallics as a source of high-current/high magnetic field superconductors". Physica C: Superconductivity 372–376 (3): 1315–1320. doi: 10.1016/S0921-4534(02)01018-3 (http:/ / dx. doi. org/ 10. 1016/ S0921-4534(02)01018-3). [57] Grunblatt, G.; Mocaer, P.; Verwaerde Ch.; Kohler, C. (2005). "A success story: LHC cable production at ALSTOM-MSA". Fusion Engineering and Design (Proceedings of the 23rd Symposium of Fusion Technology) 75–79: 1–5. doi: 10.1016/j.fusengdes.2005.06.216 (http:/ / dx. doi. org/ 10. 1016/ j. fusengdes. 2005. 06. 216). [58] Grill, Robert; Gnadenberge, Alfred (2006). "Niobium as mint metal: Production–properties–processing". International Journal of Refractory Metals and Hard Materials 24 (4): 275–282. doi: 10.1016/j.ijrmhm.2005.10.008 (http:/ / dx. doi. org/ 10. 1016/ j. ijrmhm. 2005. 10. 008). [59] " 25 Euro - 150 Years Semmering Alpine Railway (2004) (http:/ / austrian-mint. at/ bimetallmuenzen?l=en& muenzeSubTypeId=113& muenzeId=217)". 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Journal of Materials Science: Materials in Medicine 15: 1073–1077. doi: 10.1023/B:JMSM.0000046388.07961.81 (http:/ / dx. doi. org/ 10. 1023/ B:JMSM. 0000046388. 07961. 81). . [65] Biason Gomes, M. A.; Onofre, S.; Juanto, S.; Bulhões, L. O. de S. (1991). "Anodization of niobium in sulphuric acid media". Journal of Applied Electrochemistry 21 (11): 1023–1026. doi: 10.1007/BF01077589 (http:/ / dx. doi. org/ 10. 1007/ BF01077589). [66] Chiou, Y. L. (1971). "A note on the thicknesses of anodized niobium oxide films". Thin Solid Films 8 (4): R37–R39. doi: 10.1016/0040-6090(71)90027-7 (http:/ / dx. doi. org/ 10. 1016/ 0040-6090(71)90027-7).
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Niobium [67] Azevedo, C. R. F.; Spera, G.; Silva, A. P. (2002). " Characterization of metallic piercings that caused adverse reactions during use (http:/ / www. springerlink. com/ content/ 575x64408lnk560j/ )". Journal of Failure Analysis and Prevention 2 (4): 47–53. doi:10.1007/BF02715453 (inactive 2008-11-12) . . [68] Henderson, Stanley Thomas; Marsden, Alfred Michael; Hewitt, Harry (1972). Lamps and Lighting. Edward Arnold Press. pp. 244–245. ISBN 0-7131-3267-1. [69] Eichelbrönner, G. (1998). "Refractory metals: crucial components for light sources" (PDF). International Journal of Refractory Metals and Hard Materials 16 (1): 5–11. doi: 10.1016/S0263-4368(98)00009-2 (http:/ / dx. doi. org/ 10. 1016/ S0263-4368(98)00009-2). [70] Michaluk, Christopher A.; Huber, Louis E.; Ford, Robert B. (2001). Minerals, Metals and Materials Society, Metals and Materials Society Minerals. ed. "Niobium and Niobium 1% Zirconium for High Pressure Sodium (HPS) Discharge Lamps". Niobium Science & Technology: Proceedings of the International Symposium Niobium 2001 (Orlando, Florida, USA) (Niobium 2001 Ltd, 2002). ISBN 9780971206809. [71] US5254836 (http:/ / patft. uspto. gov/ netacgi/ nph-Parser?patentnumber=5254836) (1993-10-19) Okada, Yuuji; Kobayashi, Toshihiko; Sasabe, Hiroshi; Aoki, Yoshimitsu; Nishizawa, Makoto; Endo, Shunji, Method of arc welding with a ferrite stainless steel welding rod. [72] Pozdeev, Y. (1991). "Reliability comparison of tantalum and niobium solid electrolytic capacitors". Quality and Reliability Engineering International 14 (2): 79–82. doi: 10.1002/(SICI)1099-1638(199803/04)14:2<79::AID-QRE163>3.0.CO;2-Y (http:/ / dx. doi. org/ 10. 1002/ (SICI)1099-1638(199803/ 04)14:2<79::AID-QRE163>3. 0. CO;2-Y). [73] Jahnke, L.P.; Frank, R.G.; Redden, T.K. (1960). " Columbium Alloys Today (http:/ / www. osti. gov/ energycitations/ product. biblio. jsp?osti_id=4183692)". Metal Progr. 77 (6): 69–74. . [74] Nikulina, A. V. (2003). "Zirconium-Niobium Alloys for Core Elements of Pressurized Water Reactors". Metal Science and Heat Treatment 45 (7–8): 287–292. doi: 10.1023/A:1027388503837 (http:/ / dx. doi. org/ 10. 1023/ A:1027388503837). [75] Lilje, L.; Kakob, E.; Kostina, D.; Matheisena, A.; Möllera, W. -D.; Procha, D.; Reschkea, D.; Saitob, K. Schmüserc, P.; Simrocka, S.; Suzukid T.; Twarowskia, K. (2004). "Achievement of 35 MV/m in the superconducting nine-cell cavities for TESLA". Nuclear Instruments and Methods in Physics Research Section A: Accelerators, Spectrometers, Detectors and Associated Equipment 524: 1–12. doi: 10.1016/j.nima.2004.01.045 (http:/ / dx. doi. org/ 10. 1016/ j. nima. 2004. 01. 045). [76] Cherednichenko, Sergey; Drakinskiy, Vladimir; Berg, Therese; Khosropanah, Pourya; Kollberg, Erik (2008). "A Hot-electron bolometer terahertz mixers for the Herschel Space Observatory". Review of Scientific Instruments 79: 0345011–03451010. doi: 10.1063/1.2890099 (http:/ / dx. doi. org/ 10. 1063/ 1. 2890099). [77] Vilaplana, J.; Romaguera, C.; Grimalt, F.; Cornellana, F. (1990). "New trends in the use of metals in jewellery". Contact Dermatitis 25 (3): 145–148. doi: 10.1111/j.1600-0536.1991.tb01819.x (http:/ / dx. doi. org/ 10. 1111/ j. 1600-0536. 1991. tb01819. x). [78] Vilaplana, J.; Romaguera, C. (1998). "New developments in jewellery and dental materials". Contact Dermatitis 39 (2): 55–57. doi: 10.1111/j.1600-0536.1998.tb05832.x (http:/ / dx. doi. org/ 10. 1111/ j. 1600-0536. 1998. tb05832. x). [79] Haley, Thomas J.; Komesu, N.; Raymond, K. (1962). "Pharmacology and toxicology of niobium chloride". Toxicology and Applied Pharmacology 4 (3): 385–392. doi: 10.1016/0041-008X(62)90048-0 (http:/ / dx. doi. org/ 10. 1016/ 0041-008X(62)90048-0). [80] Downs, William L.; Scott, James K.; Yuile, Charles L.; Caruso, Frank S.; Wong, Lawrence C. K. (1965). " The Toxicity of Niobium Salts (http:/ / www. informaworld. com/ smpp/ content~content=a789022664~db=all)". American Industrial Hygiene Association Journal 26 (4): 337–346. doi: 10.1080/00028896509342740 (http:/ / dx. doi. org/ 10. 1080/ 00028896509342740). . [81] Schroeder, Henry A.; Mitchener, Marian; Nason, Alexis P. (01 Jan 1970). " Zirconium, Niobium, Antimony, Vanadium and Lead in Rats: Life term studies (http:/ / jn. nutrition. org/ cgi/ content/ abstract/ 100/ 1/ 59)". Journal of Nutrition 100 (1): 59–68. PMID 5412131. . [82] [83] [84] [85]
http:/ / periodic. lanl. gov/ elements/ 41. html http:/ / www. webelements. com/ webelements/ elements/ text/ Nb/ index. html http:/ / www. tanb. org/ http:/ / www. symmetrymag. org/ cms/ ?pid=1000173
15
Article Sources and Contributors
Article Sources and Contributors Niobium Source: http://en.wikipedia.org/w/index.php?oldid=306587341 Contributors: .adm.bmb., 209.75.42.xxx, Ahoerstemeier, AlanBurlison, AlimanRuna, Anclation, Andres, Archimerged, Arkuat, Art LaPella, Art-top, Attilios, Awadewit, Axiosaurus, Axl, BRG, Baccyak4H, Benbest, Bender235, BillFlis, Blackeagle, BlueEarth, BorgQueen, Borislav Dopudja, Breno, Bridesmill, Bryan Derksen, C+C, CYD, Calvin 1998, Carcharoth, CardinalDan, Carnildo, Casliber, Conversion script, Crag, Cryptic C62, Crystal whacker, Daisystanton, Dana boomer, Darrien, Dave1959, David Latapie, Dblandford, Delta G, Deor, DerHexer, Derek Ross, Dhc5007, Dispenser, Doktorbuk, Donarreiskoffer, Download, DrKiernan, DragonflySixtyseven, Drift3r insid3, Dschwen, Dwmyers, Dysprosia, EPO, EchetusXe, Edgar181, Ef3, El C, Elassint, Elchupachipmunk, Emperorbma, Epbr123, Everyguy, Exukvera, FallingRain123, Femto, Francs2000, Fresheneesz, Fusionmix, Gene Nygaard, GenestealerUK, Gh, Giftlite, Globe Collector, Grendelkhan, Gurch, Hadal, Hak-kâ-ngìn, Hallpriest9, Hannibal, Helge Skjeveland, Heron, Icairns, IceDragon64, Ideyal, Igodard, Itub, Ixfd64, J G Campbell, JFreeman, JWB, Jaraalbe, Jared81, Jarry1250, Jennavecia, Jimius, Jlin, Jmclark911, Joanjoc, John, Jonnabuz, Jose77, Kaszeta, Keenan Pepper, Kwamikagami, LA2, Lesgles, Lightmouse, LorenzoB, LouScheffer, LuigiManiac, MER-C, Makaristos, Mataaxa, Materialscientist, Mattisse, Mav, Mgimpel, Michael Devore, Michael Zimmermann, Miguel.mateo, MilFlyboy, Mild Bill Hiccup, Million Moments, Minesweeper, Mm40, Moeron, Mpatel, Ndsg, Nergaal, Nessalc, NewEnglandYankee, Nickj, Nihiltres, NinjaCharlie, Nkot, Noclevername, Nubiatech, Ocee, Olof, Oxymoron83, ParticleMan, Permax, Peter Delmonte, Peter1946, Petergans, Peytonbland, Physchim62, Physicsdavid, Piledhigheranddeeper, PlatinumX, Plexust, Poolkris, Pras, ProtoOfSnagem, Pyrope, RJHall, RTC, Rcawsey, Reinderien, Remember, Rexparry sydney, Reyk, Reywas92, Richard n, Rjwilmsi, Roberta F., Rod57, Rolinator, Romanm, Roybb95, Rreagan007, Ruslik0, SamRushing, SandyGeorgia, Saperaud, Savant13, Schneelocke, Schwanzfixierter Zuchthengst, Sengkang, Sfuerst, Shell Kinney, Simpsons contributor, Sjc, Sl, Sleepy phil, Sodium, SpikeMolec, Stevertigo, Stifynsemons, Stone, StrontriumDogs, Sunborn, Tagishsimon, Tetracube, The Vigilante Editor, TheNewPhobia, Theresa knott, Thingg, Thricecube, Tide rolls, Titus III, Tjfulopp, TonyGerillo, TopHyatt, Tpbradbury, Triwbe, Tsorro, Tsunexus, Vary, Visor, Vsmith, Walkerma, Watch37264, Wavelength, Wiki Raja, Wiki'sajoke, WolfmanSF, Woohookitty, Yyy, ZX81, 240 anonymous edits
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16
Molybdenum
1
Molybdenum 42
niobium ← molybdenum → technetium
Cr ↑
Mo ↓
W Periodic Table - Extended Periodic Table
General Name, Symbol, Number
molybdenum, Mo, 42
Element category
transition metals
Group, Period, Block
6, 5, d
Appearance
gray metallic
Standard atomic weight
95.94(2) g·mol
Electron configuration
[Kr] 4d 5s
−1
5
Electrons per shell
1
2, 8, 18, 13, 1 Physical properties
Phase
solid
Density (near r.t.)
10.28 g·cm−3
Liquid density at m.p.
9.33 g·cm−3
Melting point
2896 K (2623 °C, 4753 °F)
Boiling point
4912 K (4639 °C, 8382 °F)
Heat of fusion
37.48 kJ·mol−1
Heat of vaporization
617 kJ·mol−1
Specific heat capacity
(25 °C) 24.06 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
2742
2994
3312
3707
4212
4879
Atomic properties
Molybdenum
2
Crystal structure
body centered cubic
Oxidation states
6, 5, 4, 3, 2, 1 , -1, -2 (strongly acidic oxide)
Electronegativity Ionization energies (more)
[1]
2.16 (Pauling scale) 1st: 684.3 kJ·mol−1 2nd: 1560 kJ·mol−1 3rd: 2618 kJ·mol−1
Atomic radius
139 pm
Covalent radius
154±5 pm Miscellaneous [2]
Magnetic ordering
paramagnetic
Electrical resistivity
(20 °C) 53.4 n Ω·m
Thermal conductivity
(300 K) 138 W·m
Thermal expansion
(25 °C) 4.8 µm·m
Speed of sound (thin rod)
(r.t.) 5400 m·s
Young's modulus
329 GPa
Shear modulus
126 GPa
Bulk modulus
230 GPa
−1
−1
−1
Poisson ratio
0.31
Mohs hardness
5.5
Vickers hardness
1530 MPa
Brinell hardness
1500 MPa
CAS registry number
·K
7439-98-7 Most-stable isotopes
−1
·K
−1
Molybdenum
3
Main article: Isotopes of molybdenum iso 92
Mo
93
Mo
NA 14.84% syn
half-life
95
15.92%
95
96
16.68%
96
97
9.55%
97
98
24.13%
98
Mo Mo
99
Mo
syn
ε
4×103 y 94
Mo
Mo
9.63%
-
93
0.436, 1.214
99
Nb
Mo is stable with 52 neutron Mo is stable with 53 neutron Mo is stable with 54 neutron Mo is stable with 55 neutron Mo is stable with 56 neutron
65.94 h
−
β
γ
100
DP
Mo is stable with 50 neutron
9.25%
Mo
DE (MeV)
92
94
Mo
DM
18
7.8×10
y
− −
β β
0.74, 0.36, 0.14 3.04
Tc
-
100
Ru
References
Molybdenum (pronounced /məˈlɪbdənəm/, from the Greek word for the metal "lead"), is a Group 6 chemical element with the symbol Mo and atomic number 42. The free element, which is a silvery metal, has the sixth-highest melting point of any element. It readily forms hard, stable carbides, and for this reason it is often used in high-strength steel alloys. Molybdenum does not occur as the free metal in nature, but rather in a variety of oxidation states in minerals. Industrially molybdenum compounds are used in high-pressure and temperature resistant greases between metals, as pigments, and catalysts. Molybdenum minerals have long been known, but the element was "discovered" (in the sense of differentiating it as a new entity from minerals salts of other metals) in 1778 by Carl Wilhelm Scheele. The metal was first isolated in 1781 by Peter Jacob Hjelm. Most of molybdenum's compounds have low water solubility, but the molybdate ion MoO2−4 is soluble, and will form if molybdenum-containing minerals are in contact with free oxygen and water. Recent theories suggest that the release of free oxygen by early life was important in removing molybdenum from minerals into a soluble form in the early oceans, where it was available to be used as a catalyst by single-celled organisms. This sequence may have been important in the history of life, because molybdenum-containing enzymes then became the most important catalysts used by some bacteria to break the bond in atmospheric molecular nitrogen, allowing biological nitrogen fixation. This, in turn allowed biologically driven nitrogen-fertilization of the oceans, and thus the development of more complex organisms. Aside from bacterial enzymes involved with nitrogen fixation, about 20 different molybdenum-containing enzymes are known today in animals. Molybdenum is a required element for life in these higher organisms, though not in all bacteria.
Molybdenum
4
Characteristics Physical Molybdenum is a transition metal with an electronegativity of 1.8 on the Pauling scale and an atomic mass of 95.94 g/mole.[3] It does not react with oxygen or water at room temperature. At elevated temperatures, molybdenum trioxide is formed in the reaction 2 Mo + 3 O2 → 2 MoO3.[4] In its pure metal form, molybdenum is silvery white with a Mohs hardness of 5.5, though it is somewhat more ductile than tungsten. It has a melting point of 2623 °C (4753 °F); of the naturally occurring elements, only tantalum, osmium, rhenium, tungsten, and carbon have higher melting points.[5] Molybdenum burns only at temperatures above 600 °C (1112 °F).[6] It has the lowest heating expansion of any commercially used metal.[7] Molybdenum has a value of approximately $65,000 per tonne as of 4 May 2007. It maintained a price at or near $10,000 per tonne from 1997 through 2002, and reached a high of $103,000 per tonne in June 2005.[8]
Occurrence The world's largest producers of molybdenum materials are the United States, Canada, Chile, Russia, and China.[7] [9]
Molybdenum output in 2005
Though molybdenum is found in such minerals as wulfenite (PbMoO4) and powellite (CaMoO4), the main commercial source of molybdenum is molybdenite (MoS2). Molybdenum is mined as a principal ore, and is also recovered as a byproduct of copper and tungsten mining.[5] Large mines in Colorado (such as the Henderson mine and the now-inactive Climax mine)[10] and in British Columbia yield molybdenite as their primary product, while many porphyry copper deposits Molybdenite on quartz such as the Bingham Canyon Mine in Utah and the Chuquicamata mine in northern Chile produce molybdenum as a byproduct of copper mining. The Knaben mine in southern Norway was opened in 1885, making it the first molybdenum mine. It remained open until 1973. Molybdenum is the 54th most abundant element in the Earth's crust, and the 25th most abundant element in the oceans, with an average of 10 parts per billion; it is 42nd most abundant element in the Universe.[6] [7] The Russian Luna 24 mission discovered a single molybdenum-bearing grain (1 × 0.6 µm) in a pyroxene fragment taken from Mare Crisium on the Moon.[11] A side product of molybdenum mining is rhenium. As it is always present in small varying quantities in molybdenite, the only commercial source for rhenium is molybdenum mines.
Molybdenum
5
Isotopes There are 35 known isotopes of molybdenum ranging in atomic mass from 83 to 117, as well as four metastable nuclear isomers. Seven isotopes occur naturally, with atomic masses of 92, 94, 95, 96, 97, 98, and 100. Of these naturally occurring isotopes, only molybdenum-92 and molybdenum-100 are unstable.[12] All unstable isotopes of molybdenum decay into isotopes of niobium, technetium, and ruthenium.[13] Molybdenum-98 is the most abundant isotope, comprising 24.14% of all molybdenum. Molybdenum-100 has a half-life of approximately 1×1019 y and undergoes double beta decay into ruthenium-100. Molybdenum isotopes with mass numbers from 111 to 117 all have half-lives of approximately 150 ns.[12] [13] As also noted below, the most common isotopic molybdenum application involves molybdenum-99, which is a fission product. It is used as a parent radioisotope to the short-lived gamma-emitting daughter radioisotope technetium-99m, a nuclear isomer which is used in various imaging applications in medicine.[14]
Compounds and chemistry Oxidation states [15] of molybdenum. −2
Na2[Mo2(CO)10]
0
Mo(CO)6
+1
Na[C6H6Mo]
+2
MoCl2
+3
Na3[Mo(CN)]6
+4
MoS2
+5
MoCl5
+6
MoF6
Molybdenum has several common oxidation states, +2, +3, +4, +5 and +6. The chemistry and the compounds show more similarity to those of tungsten than that of chromium. An example is the instability of molybdenum(III) and tungsten(III) compounds compared to the stability of the chromium(III) compounds. The highest oxidation state is common in the molybdenum(VI) oxide MoO3 while the normal sulfur compound is molybdenum disulfide MoS2.
Molybdenum
Keggin structure of the phosphomolybdate anion (P[Mo12O40]3-), an example of a polyoxometalate
6 Molybdenum(VI) oxide is soluble in alkaline water, forming molybdates (MoO2−4). Molybdates are weaker oxidants than chromates, but they show a similar tendency to form complex oxyanions by condensation at lower pH values, such as [Mo7O24]6- and [Mo8O26]4-. Polymolybdates can incorporate other ions into their structure, forming polyoxometalates.[16] The dark blue phosphorus-containing heteropolymolybdate 3P[Mo12O40] is used for the spectroscopic detection of phosphorus. Molybdenum has a broad range of oxidation states, several of which are demonstrated by the various compounds of molybdenum and chlorine: • Molybdenum(II) chloride MoCl2 (yellow solid) • Molybdenum(III) chloride MoCl3 (dark red solid) • Molybdenum(V) chloride MoCl5 (dark green solid)
• Molybdenum(VI) chloride MoCl6 (brown solid) The structure of the MoCl2 is composed of Mo6Cl84+ clusters with four chloride ions to compensate the charge. Like chromium and some other transition metals molybdenum is able to form quadruple bonds. Mo2(CH3COO)4 is an example for a quadruple bond. This compound can be transformed into the chlorine compound Mo2Cl84-. The oxidation state 0 is possible with carbon monoxide as ligand, such as in molybdenum hexacarbonyl, Mo(CO)6.
History Molybdenite (from the Ancient Greek Μόλυβδος molybdos, meaning lead),[5] the principal ore from which molybdenum is now extracted, was previously known as molybdena. Molybdena was confused with and often implemented as though it were graphite. Even when the two ores were distinguishable, molybdena was thought to be a lead ore.[7] In 1754, Bengt Qvist examined the mineral and determined that it did not contain lead.[17] It was not until 1778 that Swedish chemist Carl Wilhelm Scheele realized molybdena was neither graphite nor lead.[18] [19] He and other chemists then correctly assumed that it was the ore of a distinct new element, named molybdenum for the mineral in which it was discovered. Peter Jacob Hjelm successfully isolated molybdenum using carbon and linseed oil in 1781.[7] [20] For a long time there was no industrial use for molybdenum. The French Schneider Electrics company produced the first steel molybdenum alloy armor plates in 1894. Until World War I most other armor factories also used molybdenum alloys. In World War I, some British tanks were protected by 75 mm (3.0 in) manganese plating, but this proved to be ineffective. The manganese plates were then replaced with 25 mm (0.98 in) molybdenum plating. These allowed for higher speed, greater maneuverability, and, despite being thinner, better protection.[7] The high demand for molybdenum in World War I and World War II and the steep decrease after the wars had a great influence on prices and production of molybdenum. In 2008 the London Metal Exchange announced that molybdenum would be be traded as a commodity on the exchange.[21]
Molybdenum
7
Production The molybdenite is first heated to a temperature of 700 °C (1292 °F) and the sulfide is oxidized into molybdenum(VI) oxide by air: 2 MoS2 + 7 O2 → 2 MoO3 + 4 SO2 The oxidized ore is then either heated to 1100 °C (2010 °F) to sublimate the oxide, or leached with ammonia which reacts with the molybdenum(VI) oxide to form water-soluble molybdates: MoO3 + 2 NH4OH → (NH4)2(MoO4) + H2O Copper, an impurity in molybdenite, is less soluble in ammonia. To completely remove it from the solution, it is precipitated with hydrogen sulfide. Pure molybdenum is produced by reduction of the oxide with hydrogen, while the molybdenum for steel production is reduced by the aluminothermic reaction with addition of iron to produce ferromolybdenum. A common form of ferromolybdenum contains 60% molybdenum.[7] [22]
Applications The ability of molybdenum to withstand extreme temperatures without significantly expanding or softening makes it useful in applications that involve intense heat, including the manufacture of aircraft parts, electrical contacts, industrial motors, and filaments.[7] [23] Molybdenum is also used in alloys for its high corrosion resistance and weldability.[6] [24] Molybdenum contributes corrosion resistance to type 316 stainless steel by 'gettering' residual carbon, preventing the formation of chromium carbide at grain boundaries. Most high-strength steel alloys are .25% to 8% molybdenum.[5] Despite being used in such small portions, more than 43 million kg of molybdenum is used as an alloying agent each year in stainless steels, tool steels, cast irons, and high-temperature superalloys.[6] Because of its lower density and more stable price, molybdenum is implemented in the place of tungsten.[6] An example is the 'M' series of high-speed steels such as M2, M4, and M42 as substitution for the 'T' series of HSS. Molybdenum can be implemented both as an alloying agent and as a flame-resistant coating for other metals. Although its melting point is 2623 °C (4753 °F), molybdenum rapidly oxidizes at temperatures above 760 °C (1400 °F), making it better-suited for use in vacuum environments.[23] Molybdenum based alloys have only limited applications. Due to the corrosion resistance against molten zinc, molybdenum and the molybdenum tungsten alloy (70%/30%) are used for piping, stirrers and pump impellers which come into contact with molten zinc.[25] Molybdenum-99 is used as a parent radioisotope to Technetium-99m, which is used in many medical procedures.
the
daughter
radioisotope
Molybdenum disulfide (MoS2) is used as a solid lubricant and an extreme pressure (EP) antiwear agent. It forms strong films on metallic surfaces, and is highly resistant to both extreme temperatures and high pressure, and for this reason, it is a common additive to extreme pressure application greases; in case of a catastrophic failure, the thin layer of molybdenum prevents metal-on-metal contact. Molybdenum trioxide (MoO3) is used as an adhesive between enamels and metals.[18] Molybdenum powder is used as a fertilizer for some plants, such as cauliflower.[6]
Molybdenum
8
Lead molybdate (Wulfenite) co-precipitated with lead chromate and lead sulfate is a bright-orange pigment used with ceramics and plastics.[26] Also used in NO, NO2, NOx analyzers in power plants for pollution controls. At 350 °C (662 °F) the element acts as a catalyst for NO2/NOx to form only NO molecules for consistent readings by infrared light.
Biological role The most important use of the molybdenum atom in living organisms is as a metal hetero-atom at the active site in certain enzymes. In nitrogen fixation in certain bacteria, the nitrogenase enzyme which is involved in the terminal step of reducing molecular nitrogen, usually contains molybdenum in the active site (though replacement of Mo with iron or vanadium is known). The structure of the catalytic center of the enzyme is similar to that in iron-sulfur proteins, it incorporates a Fe4S3 and MoFe3S3 cluster.[27] In March 2008, researchers reported that they had found strong evidence for the hypothesis that a scarcity of molybdenum in the Earth's early oceans was a limiting factor in the further evolution of eukaryotic life (which includes all plants and animals) as eukaryotes cannot fix nitrogen and must acquire it from prokaryotic bacteria.[28] [29] [30] The scarcity of molybdenum resulted from the relative lack of oxygen in the early ocean. Oxygen dissolved in seawater is the primary mechanism for dissolving molybdenum from minerals on the sea bottom. Though molybdenum forms compounds with various organic molecules, including carbohydrates and amino acids, it is transported throughout the human body as MoO2−4.[31] Molybdenum is present in approximately 20 enzymes in animals, including aldehyde oxidase, sulfite oxidase and xanthine oxidase.[7] In some animals, the oxidation of xanthine to uric acid, a process of purine catabolism, is catalyzed by xanthine Molybdenum containing cofactor molybdopterin oxidase, a molybdenum-containing enzyme. The activity of xanthine oxidase is directly proportional to the amount of molybdenum in the body. However, an extremely high concentration of molybdenum reverses the trend, and can act as an inhibitor in both purine catabolism and other processes. Molybdenum concentrations also affect protein synthesis, metabolism, and growth.[31] These enzymes in plants and animals catalyze the reaction of oxygen in small molecules, as part of the regulation of nitrogen-, sulfur- and carbon cycles.[32] Human body contains about 0.07 mg of molybdenum per kilogram.[33] It occurs in higher concentrations in the liver and kidneys, and in lower concentrations in the vertebrae.[6] Molybdenum is also present within human tooth enamel and may help prevent the decaying thereof.[34] Pork, lamb, and beef liver each have approximately 1.5 parts per million of molybdenum. Other significant dietary sources include green beans, eggs, sunflower seeds, wheat flour, lentils, and cereal grain.[7] The average daily intake of molybdenum varies between 120 µg and 240 µg but strongly depending on the molybdenum content of the food.[35] Acute toxicity has not been seen in humans, and the toxicity depends strongly on the chemical state. Studies on rats show a median lethal dose (LD50) as low as 180 mg/kg for some Mo compounds.[36] Although
Molybdenum
9
human toxicity data is unavailable, animal studies have shown that chronic ingestion of more than 10 mg/kg of molybdenum can cause diarrhea, growth retardation, sterility, low birth weight, and gout, as well as affecting the lungs, kidneys, and liver. Molybdenum deficiency is not usually seen in healthy people.[37] Sodium tungstate is a competitive inhibitor of molybdenum. Dietary tungsten reduces the concentration of molybdenum in tissues.[6] Dietary deficiency in molybdenum from low soil concentration has been associated with increased rates of esophageal cancer in a geographical band from [[northern China to Iran.[38] [39] Compared to the United States, which has a greater supply of molybdenum in the soil, those living in these areas have a 16 times greater risk for esophageal squamous cell carcinoma.
Copper-molybdenum antagonism High levels of molybdenum can interfere with the body's uptake of copper, producing copper deficiency. Molybdenum prevents plasma proteins from binding to copper, and it also increases the amount of copper that is excreted in urine. Ruminants that consume high amounts of molybdenum develop symptoms including diarrhea, stunted growth, anemia, and achromotrichia (loss of hair pigment). These symptoms can be alleviated by the administration of more copper into the system, both in dietary form and by injection.[40] The condition can be aggravated by excess sulfur.[6]
Precautions Molybdenum dusts and fumes, as can be generated by mining or metalworking, can be toxic, especially if ingested (including dust trapped in the sinuses and later swallowed).[36] Low levels of prolonged exposure can cause irritation to the eyes and skin. The direct inhalation or ingestion of molybdenum and its oxides should also be avoided.[41] [42] OSHA regulations specify the maximum permissible molybdenum exposure in an 8-hour day to be 5 mg/m³. Chronic exposure to 60 to 600 mg Mo/m³ can cause symptoms including fatigue, headaches, and joint pains.[43]
External links • WebElements.com [44] — Molybdenum • International Molybdenum Association
[45]
— Main page
References [1] " Molybdenum: molybdenum(I) fluoride compound data (http:/ / openmopac. net/ data_normal/ molybdenum(i) fluoride_jmol. html)". OpenMOPAC.net. . Retrieved 2007-12-10. [2] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [3] " Properties of Molybdenum (http:/ / www. qivx. com/ ispt/ elements/ ptw_042. php)". Integral Scientist Periodic Table. Qivx, Inc.. 2003. . Retrieved 2007-06-10. [4] Winter, Mark. " Chemistry (http:/ / www. webelements. com/ webelements/ elements/ text/ Mo/ chem. html)". Molybdenum. The University of Sheffield. . Retrieved 2007-06-10. [5] Lide, David R., ed. (1994), "Molybdenum", CRC Handbook of Chemistry and Physics, 4, Chemical Rubber Publishing Company, p. 18, ISBN 0849304741 [6] Considine, Glenn D., ed (2005). "Molybdenum". Van Nostrand's Encyclopedia of Chemistry. New York: Wiley-Interscience. pp. 1038–1040. ISBN 9780471615255.
Molybdenum [7] Emsley, John (2001). Nature's Building Blocks (http:/ / books. google. com/ books?id=j-Xu07p3cKwC& pg=PA265). Oxford: Oxford University Press. pp. 262–266. ISBN 0198503415. . [8] " Dynamic Prices and Charts for Molybdenum (http:/ / www. infomine. com/ investment/ metalschart. asp?c=molybdenum& u=mt& submit1=Display+ Chart& x=usd& r=15y)". InfoMine Inc.. 2007. . Retrieved 2007-05-07. [9] Lide, David R., ed (2006). CRC Handbook of Chemistry and Physics. 4. Chemical Rubber Publishing Company. pp. 22–23. ISBN 0849304873. [10] Coffman, Paul B. (1937). "The Rise of a New Metal: The Growth and Success of the Climax Molybdenum Company". The Journal of Business of the University of Chicago 10: 30. doi: 10.1086/232443 (http:/ / dx. doi. org/ 10. 1086/ 232443). [11] " New mineral names (http:/ / www. minsocam. org/ msa/ AmMin/ TOC/ Abstracts/ 2002_Abstracts/ Jan02_Abstracts/ Jambor_p181_02. pdf)". American Mineralogist 87: 181. 2002. . [12] Audi, Georges (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3–128. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [13] Lide, David R., ed. (2006), CRC Handbook of Chemistry and Physics, 11, CRC, pp. 87–88, ISBN 0849304873 [14] Armstrong, John T. (2003). " Technetium (http:/ / pubs. acs. org/ cen/ 80th/ technetium. html)". Chemical & Engineering News. . Retrieved 2009-07-07. [15] Schmidt, Max (1968). "VI. Nebengruppe" (in German). Anorganische Chemie II.. Wissenschaftsverlag. pp. 119–127. [16] Pope, Michael T.; Müller, Achim (1997). "Polyoxometalate Chemistry: An Old Field with New Dimensions in Several Disciplines". Angewandte Chemie International Edition 30: 34. doi: 10.1002/anie.199100341 (http:/ / dx. doi. org/ 10. 1002/ anie. 199100341). [17] Van der Krogt, Peter (2006-01-10). " Molybdenum (http:/ / www. vanderkroft. net/ elements/ elem/ mo. html)". Elementymology & Elements Multidict. . Retrieved 2007-05-20. [18] Gagnon, Steve. " Molybdenum (http:/ / education. jlab. org/ itselemental/ ele042. html)". Jefferson Science Associates, LLC. . Retrieved 2007-05-06. [19] C. W. K. Scheele (1779). " Versuche mit Wasserbley;Molybdaena (http:/ / resolver. sub. uni-goettingen. de/ purl?PPN324352840_0040)". Svenska vetensk. Academ. Handlingar 40: 238. . [20] P. J. Hjelm (1788). " Versuche mit Molybdäna, und Reduction der selben Erde (http:/ / resolver. sub. uni-goettingen. de/ purl?PPN324352840_0009_02_NS)". Svenska vetensk. Academ. Handlingar 49: 268. . [21] " LME to launch minor metals contracts in H2 2009 (http:/ / lme. com/ 6241. asp)". London Metal Exchange. 4 September 2008. . Retrieved 28 July 2009. [22] Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils; (1985) (in German). Lehrbuch der Anorganischen Chemie (91–100 ed.). Walter de Gruyter. pp. 1056–1057. ISBN 3-11-007511-3. [23] " Molybdenum (http:/ / www. azom. com/ details. asp?ArticleID=616)". AZoM.com Pty. Limited. 2007. . Retrieved 2007-05-06. [24] " Molybdenum Statistics and Information (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ molybdenum/ )". U.S. Geological Survey. 2007-05-10. . Retrieved 2007-05-10. [25] Cubberly, W. H.; Bakerjian, Ramon (1989). Tool and manufacturing engineers handbook (http:/ / books. google. de/ books?id=NRXnXmFRjWYC& pg=PT421). Society of Manufacturing Engineers. p. 421. ISBN 9780872633513. . [26] International Molybdenum Association, http:/ / www. moly. imoa. info [27] Dos Santos, Patricia C.; Dean, Dennis R. (2008). "A newly discovered role for iron-sulfur clusters". PNAS 105: 11589. doi: 10.1073/pnas.0805713105 (http:/ / dx. doi. org/ 10. 1073/ pnas. 0805713105). [28] Scott, C.; et al. (2008). "Tracing the stepwise oxygenation of the Proterozoic ocean". Nature 452 (7186): 456–460. doi: 10.1038/nature06811 (http:/ / dx. doi. org/ 10. 1038/ nature06811). [29] " International team of scientists discover clue to delay of life on Earth (http:/ / www. eurekalert. org/ pub_releases/ 2008-03/ asu-ito032508. php)". Eurekalert.org. . Retrieved 2008-10-25. [30] " Scientists uncover the source of an almost 2 billion year delay in animal evolution (http:/ / www. eurekalert. org/ pub_releases/ 2008-03/ nu-sut032508. php)". Eurekalert.org. . Retrieved 2008-10-25. [31] Mitchell, Phillip C. H. (2003). " Overview of Environment Database (http:/ / www. hse. imoa. info/ Default. asp?page=110)". International Molybdenum Association. . Retrieved 2007-05-05. [32] Kisker, C.; Schindelin, H.; Baas, D.; Rétey, J.; Meckenstock, R.U.; Kroneck, P.M.H. (1999). "A structural comparison of molybdenum cofactor-containing enzymes". FEMS Microbiol. Rev. 22: 503. doi: 10.1111/j.1574-6976.1998.tb00384.x (http:/ / dx. doi. org/ 10. 1111/ j. 1574-6976. 1998. tb00384. x). PMID 9990727. [33] Arnold F. Holleman, Egon Wiberg (2001). Inorganic chemistry (http:/ / books. google. com/ books?id=vEwj1WZKThEC& pg=PA1384). Academic Press. p. 1384. ISBN 0123526515. .
10
Molybdenum [34] Ismail, Mumtaz. " Dental Problems and Diet (http:/ / www. bawarchi. com/ health/ dental. html)". Health and Nutrition. Bawarchi. . Retrieved 2007-05-19. [35] Coughlan, M. P. (1983). "The role of molybdenum in human biology". Journal of Inherited Metabolic Disease 6: 70. doi: 10.1007/BF01811327 (http:/ / dx. doi. org/ 10. 1007/ BF01811327). [36] " Risk Assessment Information System: Toxicity Summary for Molybdenum (http:/ / rais. ornl. gov/ tox/ profiles/ molybdenum_f_V1. shtml)". Oak Ridge National Laboratory. . [37] " Nutrient Reference Values for Australia (http:/ / www. nrv. gov. au/ Nutrients. aspx?code=71128006)". National Medical and Health Research Council (Australia). . Retrieved 2008-04-23. [38] Yang, Chung S. (1980). " Research on Esophageal Cancer in China: a Review (http:/ / cancerres. aacrjournals. org/ cgi/ reprint/ 40/ 8_Part_1/ 2633. pdf)". Cancer Research 40: 2633. . [39] Nouri, Mohsen; Chalian, Hamid; Bahman, Atiyeh; Mollahajian, Hamid; Ahmadi-Faghih, Mohammadamin; Fakheri, Hafez; Soroush, Ahmadreza (2008). " Nail Molybdenum and Zinc Contents in Populations with Low and Moderate Incidence of Esophageal Cancer (http:/ / www. ams. ac. ir/ AIM/ 08114/ 0010. pdf)". Archives of Iranian Medicine 11: 392. . [40] Suttle, N. F. (December 1974). "Recent studies of the copper-molybdenum antagonism". Proceedings of the Nutrition Society (CABI Publishing) 33 (3): 299–305. doi: 10.1079/PNS19740053 (http:/ / dx. doi. org/ 10. 1079/ PNS19740053). [41] " Material Safety Data Sheet - Molybdenum (http:/ / www. rembar. com/ MSDSmo. htm)". The REMBAR Company, Inc.. 2000-09-19. . Retrieved 2007-05-13. [42] " Material Safety Data Sheet - Molybdenum Powder (http:/ / asp. cerac. com/ CatalogNet/ default. aspx?p=msdsFile& msds=m000121. htm)". CERAC, Inc.. 1994-02-23. . Retrieved 2007-10-19. [43] " NIOSH Documentation for ILDHs Molybdenum (http:/ / www. cdc. gov/ niosh/ idlh/ moly-mo. html)". National Institute for Occupational Safety and Health. 1996-08-16. . Retrieved 2007-05-31. [44] http:/ / www. webelements. com/ molybdenum/ [45] http:/ / www. imoa. info/ index. html
11
Article Sources and Contributors
Article Sources and Contributors Molybdenum Source: http://en.wikipedia.org/w/index.php?oldid=306351961 Contributors: A2Kafir, AdjustShift, Ahoerstemeier, [email protected], AlimanRuna, Anclation, Andres, Andrewa, Andrewpmk, Animum, Anwar saadat, Archimerged, Attilios, Auric, Avidallred, Baccyak4H, Bakabaka, Bduke, Beetstra, Bender235, BigDunc, Biscuittin, Bkell, BlueEarth, Bobo192, BorgQueen, Brandmeister, [email protected], BrianKnez, Bryan Derksen, CDrecche, CYD, Camembert, Carnildo, Chekaz, Chloh, Chowbok, Cometstyles, Conversion script, Coppertwig, Corixidae, Crazynas, Cryptic C62, DHollerman, Darrien, David Latapie, David R. Ingham, Delta G, Depakote, Dgrant, Dismas, Dividing, DocWatson42, Doulos Christos, Dpasek, Dratman, Drini, Drphilharmonic, Dwmyers, EPO, Edgar181, Edivorce, Eldin raigmore, Emperorbma, Epbr123, Erik Zachte, Excirial, Femto, Fil Defiarse, Flowanda, Fvw, Gaius Cornelius, Gbr3, Giftlite, Giraffedata, Gpvos, GraemeL, GregorB, Grendelkhan, Hadal, Hadlock, Hak-kâ-ngìn, Hallpriest9, Harish2k1vet, HoodedMan, Hydrogen Iodide, Icairns, Imperator3733, Iridescent, Irishguy, J.delanoy, JDG, Jaganath, Jaraalbe, Jellis316, Joanjoc, John, JohnI, Jose77, Jrockley, Juliancolton, Justing magpie, Kgolp76, KnowledgeOfSelf, Kumorifox, Kurykh, Kwamikagami, LA2, Ledelste, Lights, Lon of Oakdale, Lst27, LuigiManiac, MONGO, Magnus Manske, Marek69, Martarius, Materialscientist, Mav, Mbecker, Mgimpel, Minesweeper, Morwen, Mtk180, Mwistey, Mygerardromance, Nergaal, Nihiltres, NinjaCharlie, Nohat, Nonagonal Spider, Numbo3, Nxva, Opelio, OwenX, Oxymoron83, Physchim62, Piperh, Plantsurfer, PlatinumX, Plazak, Plexust, PoliteCarbide, Poolkris, Pras, Psicorps, Quiddity, RTC, Rajah9, RegaL the Proofreader, Regibox, Remember, Renesis, RexNL, Rifleman 82, Rjwilmsi, Rmrfstar, RobertAustin, Roberta F., Rolinator, Romanm, Roux, Roybb95, Sannse, Saperaud, Sbharris, Schneelocke, Search4Lancer, Sengkang, Shafei, Silsor, Skypaint, Sl, Snowmanmelting, Sonett72, SpookyMulder, Squids and Chips, Stifynsemons, Stone, Suisui, Sunborn, TDogg310, THEN WHO WAS PHONE?, Tagishsimon, Tellyaddict, Tempodivalse, Tetracube, The Librarian, TheGWO, Think8359, Thricecube, Tmangray, Tom harrison, TonyGerillo, Tranletuhan, Ttwaring, UmbraNecat, UnitedStatesian, VASANTH S.N., Versus22, Vista4u2, Vsmith, Warut, WhatamIdoing, Whosasking, Wprlh, Xoder, Yekrats, Yellier, Yyy, Zfr, Zinnmann, Zzuuzz, 425 anonymous edits
Image Sources, Licenses and Contributors image:Mo-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Mo-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Joanjoc, Mav, Paddy, Saperaud Image:Molybdaen 1.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Molybdaen_1.jpg License: Public Domain Contributors: http://de.wikipedia.org/w/index.php?title=Benutzer:Tomihahndorf Image:2005molybdenum (mined).PNG Source: http://en.wikipedia.org/w/index.php?title=File:2005molybdenum_(mined).PNG License: Public Domain Contributors: User:Anwar_saadat/bubble_maps_(FAQ) File:Molly Hill molybdenite.JPG Source: http://en.wikipedia.org/w/index.php?title=File:Molly_Hill_molybdenite.JPG License: unknown Contributors: User:Pyrope File:Phosphotungstate-3D-polyhedra.png Source: http://en.wikipedia.org/w/index.php?title=File:Phosphotungstate-3D-polyhedra.png License: Public Domain Contributors: Benjah-bmm27, 1 anonymous edits Image:Molybdopterin.svg Source: http://en.wikipedia.org/w/index.php?title=File:Molybdopterin.svg License: Public Domain Contributors: User:Fvasconcellos
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12
Technetium
1
Technetium 43
molybdenum ← technetium → ruthenium
Mn ↑
Tc ↓
Re Periodic Table - Extended Periodic Table
General Name, Symbol, Number
technetium, Tc, 43
Element category
transition metals
Group, Period, Block
7, 5, d
Appearance
silvery gray metal
Standard atomic weight
[98](0) g·mol
Electron configuration
[Kr] 4d 5s
−1
5
Electrons per shell
2
2, 8, 18, 13, 2 Physical properties
Phase
solid
Density (near r.t.)
11 g·cm−3
Melting point
2430 K (2157 °C, 3915 °F)
Boiling point
4538 K (4265 °C, 7709 °F)
Heat of fusion
33.29 kJ·mol−1
Heat of vaporization
585.2 kJ·mol−1
Specific heat capacity
(25 °C) 24.27 J·mol−1·K−1 Vapor pressure (extrapolated)
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
2727
2998
3324
3726
4234
4894
Atomic properties Crystal structure
hexagonal
Technetium
2 [1]
Oxidation states
[2]
7, 6, 5, 4, 3, 2, 1 , -1, -3 (strongly acidic oxide)
Electronegativity
1.9 (Pauling scale)
Electron affinity
-53 kJ/mol
Ionization energies
1st: 702 kJ/mol 2nd: 1470 kJ/mol 3rd: 2850 kJ/mol
Atomic radius
136 pm
Covalent radius
147±7 pm Miscellaneous
Magnetic ordering
Paramagnetic
Thermal conductivity
(300 K) 50.6 W·m
−1
CAS registry number
−1
·K
7440-26-8 Most-stable isotopes
Main article: Isotopes of technetium iso 95m
Tc
96
Tc
NA syn
syn
half-life 61 d
4.3 d
DM ε
DE (MeV) -
γ
0.204, 0.582, 0.835
IT
0.0389, e
ε γ
0.778, 0.849, 0.812
DP 95
Mo
-
95
Tc
96
Mo
-
syn
2.6×106 y
ε
-
97
97m
syn
90 d
IT
0.965, e
97
98
syn
4.2×106 y
0.4
98
97
Tc Tc
Tc
β− γ
99
Tc
99m
Tc
trace
2.111×105 y
trace
6.01 h
0.745, 0.652
Mo Tc Ru
-
0.294
99
IT
0.142, 0.002
99
γ
0.140
β−
Ru Tc
-
References
Technetium (pronounced /tɛkˈniːʃɪəm/) is the lightest chemical element with no stable isotope, and therefore the lightest radioactive element (but not the lightest radioactive isotope, which is tritium). It is a synthetic element with the atomic number 43 and is given the symbol Tc. The chemical properties of this silvery grey, crystalline transition metal are intermediate between rhenium and manganese. Its short-lived gamma-emitting nuclear isomer 99mTc (technetium-99m) is used in nuclear medicine for a wide variety of diagnostic
Technetium tests. 99Tc is used as a gamma ray-free source of beta particles. The pertechnetate ion (TcO4-) has been suggested as a strong anodic corrosion inhibitor for mild steel in closed cooling systems. Before the element was discovered, many of the properties of element 43 were predicted by Dmitri Mendeleev. Mendeleev noted a gap in his periodic table and called the element ekamanganese (Em). In 1937 its isotope 97Tc became the first predominantly artificial element to be produced, hence its name (from the Greek τεχνητός, meaning "artificial"). Most technetium produced on Earth is a by-product of fission of uranium-235 in nuclear reactors and is extracted from nuclear fuel rods. No isotope of technetium has a half-life longer than 4.2 million years (98Tc), so its detection in red giants in 1952 helped bolster the theory that stars can produce heavier elements. On Earth, technetium occurs in trace but measurable quantities as a product of spontaneous fission in uranium ore or by neutron capture in molybdenum ores.
Characteristics Technetium is a silvery-grey radioactive metal with an appearance similar to platinum. However, it is commonly obtained as a grey powder. Its position in the group 7 of the periodic table is between rhenium and manganese and as predicted by the periodic law its properties are intermediate between those two elements. Technetium, together with promethium, is unusual among the lighter elements in that it has no stable isotopes. Technetium is therefore extremely rare on Earth. Technetium plays no natural biological role and is not normally found in the human body.[3] The metal form of technetium slowly tarnishes in moist air. Its oxides are TcO2 and Tc2O7. Under oxidizing conditions technetium (VII) will exist as the pertechnetate ion, TcO4-. Common oxidation states of technetium include 0, +2, +4, +5, +6 and +7. Technetium will burn in oxygen when in powder form.[4] It dissolves in aqua regia, nitric acid, and concentrated sulfuric acid, but it is not soluble in hydrochloric acid of any strength.[3] It has characteristic spectral lines at 363 nm, 403 nm, 410 nm, 426 nm, 430 nm, and 485 nm.[5] The metal form is slightly paramagnetic, meaning its magnetic dipoles align with external magnetic fields even though technetium is not normally magnetic.[4] The crystal structure of the metal is hexagonal close-packed. Pure metallic single-crystal technetium becomes a type II superconductor at 7.46 K; irregular crystals and trace impurities raise this temperature to 11.2 K for 99.9% pure technetium powder.[6] Below this temperature technetium has a very high magnetic penetration depth, the largest among the elements apart from niobium.[7] Technetium is produced in quantity by nuclear fission, and spreads more readily than many radionuclides. In spite of the importance of understanding its toxicity in animals and humans, experimental evidence is scant. It appears to have low chemical toxicity. Its radiological toxicity (per unit of mass) is a function of compound, type of radiation for the isotope in question, and the isotope's half-life. Technetium-99m is particularly attractive for medical applications, as the radiation from this isotope is a gamma ray with the same wavelength as X-rays used for common medical diagnostic X-ray applications, giving it adequate penetration while causing minimal damage for a gamma photon. This, plus the extremely short half-life of this metastable nuclear isomer, followed by the relatively long half-life of the daughter isotope Tc-99 which allows it to be eliminated from the body before it decays. This leads to a relatively low dose of administered radiation in biologically
3
Technetium
4
dose-equivalent amounts (sieverts) for a typical Tc-99m based nuclear scan (see more on this subject below).[8] All isotopes of technetium must be handled carefully. The most common isotope, technetium-99, is a weak beta emitter; such radiation is stopped by the walls of laboratory glassware. Soft X-rays are emitted when the beta particles are stopped, but as long as the body is kept more than 30 cm away these should pose no problem. The primary hazard when working with technetium is inhalation of dust; such radioactive contamination in the lungs can pose a significant cancer risk. For most work, careful handling in a fume hood is sufficient; a glove box is not needed.[9]
History Search for element 43 For a number of years there was a gap in the periodic table between molybdenum (element 42) and ruthenium (element 44). Many early researchers were eager to be the first to discover and name the missing element; its location in the table suggested that it should be easier to find than other undiscovered elements. It was first thought to have been found in platinum ores in 1828. It was given the name polinium but it turned out to be impure iridium. Then in 1846 the element ilmenium was claimed to have been discovered but was determined to be impure niobium. This mistake was repeated in 1847 with the "discovery" of pelopium.[10] Dmitri Mendeleev predicted that this missing element, as part of other predictions, would be chemically similar to manganese and gave it the name ekamanganese.
Dmitri Mendeleev predicted technetium's properties before it was discovered.
In 1877, the Russian chemist Serge Kern reported discovering the missing element in platinum ore. Kern named what he thought was the new element davyum, after the noted English chemist Sir Humphry Davy, but it was determined to be a mixture of iridium, rhodium and iron. Another candidate, lucium, followed in 1896 but it was determined to be yttrium. Then in 1908 the Japanese chemist Masataka Ogawa found evidence in the mineral thorianite which he thought indicated the presence of element 43. Ogawa named the element nipponium, after Japan (which is Nippon in Japanese). In 2004 H. K Yoshihara utilized "a record of X-ray spectrum of Ogawa's nipponium sample from thorianite [which] was contained in a photographic plate preserved by his family. The spectrum was read and indicated the absence of the element 43 and the presence of the element 75 (rhenium)."[11] German chemists Walter Noddack, Otto Berg and Ida Tacke (later Mrs. Noddack) reported the discovery of element 75 and element 43 in 1925 and named element 43 masurium (after Masuria in eastern Prussia, now in Poland, the region where Walter Noddack's family originated).[12] The group bombarded columbite with a beam of electrons and deduced element 43 was present by examining X-ray diffraction spectrograms. The wavelength of the X-rays produced is related to the atomic number by a formula derived by Henry Moseley in 1913. The team claimed to detect a faint X-ray signal at a wavelength produced by element 43. Contemporary experimenters could not replicate the discovery, and in fact it
Technetium was dismissed as an error for many years.[13]
5 [14]
In 1998 John T. Armstrong of the National Institute of Standards and Technology ran "computer simulations" of the 1925 experiments and obtained results quite close to those reported by the Noddack team. He claimed that this was further supported by work published by David Curtis of the Los Alamos National Laboratory measuring the (tiny) natural occurrence of technetium.[13] [15] However, the Noddack's experimental results have never been reproduced, and they were unable to isolate any element 43. Debate still exists as to whether the 1925 team actually did discover element 43.
Official discovery and later history Discovery of element 43 was finally confirmed in a 1937 experiment at the University of Palermo in Sicily conducted by Carlo Perrier and Emilio Segrè. In the summer of 1936 Segrè and his wife visited the United States, first New York at Columbia University, where he had spent time the previous summer, and then Berkeley at Ernest O. Lawrence's Radiation Laboratory. He persuaded cyclotron inventor Lawrence to let him take back some discarded cyclotron parts that had become radioactive. In early 1937 Lawrence mailed him a molybdenum foil that had been part of the deflector in the cyclotron. Segrè enlisted his experienced chemist colleague Perrier to attempt to prove through comparative chemistry that the molybdenum activity was indeed Z = 43, an element not existent in nature because of its instability against nuclear decay. With considerable difficulty they finally succeeded in isolating three distinct decay periods (90, 80, and 50 days) that eventually turned out to be two isotopes, 95Tc and 97Tc, of technetium, the name given later by Perrier and Segrè to the first man-made element.[16] University of Palermo officials wanted them to name their discovery panormium, after the Latin name for Palermo, Panormus. The researchers instead named element 43 after the Greek word τεχνητός, meaning "artificial", since it was the first element to be artificially produced.[10] [12] Segrè returned to Berkeley and immediately sought out Glenn T. Seaborg. They isolated the technetium-99m isotope which is now used in some 10,000,000 medical diagnostic procedures annually.[17] In 1952 astronomer Paul W. Merrill in California detected the spectral signature of technetium (in particular, light at 403.1 nm, 423.8 nm, 426.8 nm, and 429.7 nm) in light from S-type red giants. These massive stars near the end of their lives were rich in this short-lived element, meaning nuclear reactions within the stars must be producing it. This evidence was used to bolster the then unproven theory that stars are where nucleosynthesis of the heavier elements occurs.[16] More recently, such observations provided evidence that elements were being formed by neutron capture in the s-process.[18] Since its discovery, there have been many searches in terrestrial materials for natural sources. In 1962, technetium-99 was isolated and identified in pitchblende from the Belgian Congo in extremely small quantities (about 0.2 ng/kg);[18] there it originates as a spontaneous fission product of uranium-238. This discovery was made by B.T. Kenna and P.K. Kuroda. There is also evidence that the Oklo natural nuclear fission reactor produced significant amounts of technetium-99, which has since decayed to ruthenium-99.[18]
Technetium
6
Occurrence and production Since technetium is unstable, only minute traces occur naturally in the Earth's crust as a spontaneous fission product of uranium. In 1999 David Curtis (see above) estimated that a kilogram of uranium contains 1 nanogram (1×10−9 g) of technetium.[16] Extraterrestrial technetium was found in some red giant stars (S-, M-, and N-types) that contain an absorption line in their spectrum indicating the presence of this element.[3] Uranium ores contain traces of Technetium
Long-lived fission products Prop: Unit:
Yield %
t½ Ma
Q* KeV
βγ *
99
Tc
.211
6.1385
126
Sn
.230
.1084
.295
.0447
93
Zr
1.53
5.4575
135
Cs
2.3
6.9110
269 β
107
Pd
6.5
1.2499
33 β
15.7
.8410
79
Se
129
I
294 β 4050
βγ
151 β 91
194
βγ
βγ
Byproduct production of Tc-99 in fission wastes In contrast with the rare natural occurrence, bulk quantities of technetium-99 are produced each year from spent nuclear fuel rods, which contain various fission products. The fission of a gram of uranium-235 in nuclear reactors yields 27 mg of 99Tc, giving technetium a fission product yield of 6.1%.[4] Other fissile isotopes also produce similar yields of technetium, e.g. 4.9% from uranium-233 or 6.21% from plutonium-239.[19] It is estimated that up to 1994, about 49,000 TBq (78 metric tons) of technetium was produced in nuclear reactors, which is by far the dominant source of terrestrial technetium.[20] However, only a fraction of the production is used commercially. As of 2005, technetium-99 in the form of ammonium pertechnate is available to holders of an ORNL permit for US$83/g plus packing charges.[3] This is higher than the spot price of platinum, which has never been worth more than $72.40 per gram. Since the yield of technetium-99 as a product of the nuclear fission of both uranium-235 and plutonium-239 is moderate, it is present in radioactive waste of fission reactors and is produced when a fission bomb is detonated. The amount of artificially produced technetium in the environment exceeds its natural occurrence by many orders of magnitude. This is
Technetium due to release by atmospheric nuclear testing along with the disposal and processing of high-level radioactive waste. Due to its high fission yield and relatively high half-life, technetium-99 is one of the main components of nuclear waste. Its decay, measured in becquerels per amount of spent fuel, is dominant at about 104 to 106 years after the [20] creation of the nuclear waste. An estimated 160 TBq (about 250 kg) of technetium-99 was released into the environment up to 1994 by atmospheric nuclear tests.[20] The amount of technetium-99 from nuclear reactors released into the environment up to 1986 is estimated to be on the order of 1000 TBq (about 1600 kg), primarily by nuclear fuel reprocessing; most of this was discharged into the sea. In recent years, reprocessing methods have improved to reduce emissions, but as of 2005 the primary release of technetium-99 into the environment is by the Sellafield plant, which released an estimated 550 TBq (about 900 kg) from 1995–1999 into the Irish Sea. From 2000 onwards the amount has been limited by regulation to 90 TBq (about 140 kg) per year.[21] As a result of nuclear fuel reprocessing, technetium has been discharged into the sea in a number of locations, and some seafood contains tiny but measurable quantities. For example, lobster from west Cumbria contains small amounts of technetium.[22] The anaerobic, spore-forming bacteria in the Clostridium genus are able to reduce Tc(VII) to Tc(IV). Clostridia bacteria play a role in reducing iron, manganese and uranium, thereby affecting these elements' solubility in soil and sediments. Their ability to reduce technetium may determine a large part of mobility of technetium in industrial wastes and other subsurface environments.[23] The long half-life of technetium-99 and its ability to form an anionic species makes it (along with 129I) a major concern when considering long-term disposal of high-level radioactive waste. In addition, many of the processes designed to remove fission products from medium-active process streams in reprocessing plants are designed to remove cationic species like caesium (e.g., 137Cs) and strontium (e.g., 90Sr). Hence the pertechnetate is able to escape through these treatment processes. Current disposal options favor burial in geologically stable rock. The primary danger with such a course is that the waste is likely to come into contact with water, which could leach radioactive contamination into the environment. The anionic pertechnetate and iodide are less able to absorb onto the surfaces of minerals so they are likely to be more mobile. By comparison plutonium, uranium, and caesium are much more able to bind to soil particles. For this reason, the environmental chemistry of technetium is an active area of research. An alternative disposal method, transmutation, has been demonstrated at CERN for technetium-99. This transmutation process is one in which the technetium (99Tc as a metal target) is bombarded with neutrons to form the short-lived 100Tc (half life = 16 seconds) which decays by beta decay to ruthenium (100Ru). If recovery of usable ruthenium is a goal, an extremely pure technetium target is needed; if small traces of the minor actinides such as americium and curium are present in the target, they are likely to undergo fission and form more fission products which increase the radioactivity of the irradiated target. The formation of 106Ru (half-life 374 days) from the fresh fission is likely to increase the activity of the final ruthenium metal, which will then require a longer cooling time after irradiation before the ruthenium can be used. The actual production of technetium-99 from spent nuclear fuel is a long process. During fuel reprocessing, it appears in the waste liquid, which is highly radioactive. After sitting
7
Technetium for several years, the radioactivity has fallen to a point where extraction of the long-lived isotopes, including technetium-99, becomes feasible. Several chemical extraction processes are used yielding technetium-99 metal of high purity.[24]
Neutron activation of molybdenum or other pure elements The metastable (a state where the nucleus is in an excited state) isotope 99mTc is produced as a fission product from the fission of uranium or plutonium in nuclear reactors. Because used fuel is allowed to stand for several years before reprocessing, all 99Mo and 99mTc will have decayed by the time that the fission products are separated from the major actinides in conventional nuclear reprocessing. The PUREX raffinate will contain a high concentration of technetium as TcO4- but almost all of this will be 99Tc. The vast majority of the 99mTc used in medical work is formed from 99Mo which is formed by the neutron activation of 98Mo. 99Mo has a half-life of 67 hours, so short-lived 99mTc (half-life: 6 hours), which results from its decay, is being constantly produced.[16] The hospital then chemically extracts the technetium from the solution by using a technetium-99m generator ("technetium cow", also occasionally called a "molybdenum cow"). By working in this way, there is no need for the complex chemical steps which would be required to separate molybdenum from a fission product mixture. This alternative method requires that an enriched uranium target be irradiated with neutrons to form 99Mo as a fission product, then separated.[25] Other technetium isotopes are not produced in significant quantities by fission; when needed, they are manufactured by neutron irradiation of parent isotopes (for example, 97Tc can be made by neutron irradiation of 96Ru).
Isotopes Technetium is one of the two elements in the first 82 that have no stable isotopes (in fact, it is the lowest-numbered element that is exclusively radioactive); the other such element is promethium. The most stable radioisotopes are 98Tc (half-life of 4.2 Ma), 97Tc (half-life: 2.6 Ma) and 99Tc (half-life: 211.1 ka).[26] Twenty-two other radioisotopes have been characterized with atomic masses ranging from 87.933 u (88Tc) to 112.931 u (113Tc). Most of these have half-lives that are less than an hour; the exceptions are 93Tc (half-life: 2.75 hours), 94Tc (half-life: 4.883 hours), 95Tc (half-life: 20 hours), and 96Tc (half-life: 4.28 days).[26] Technetium also has numerous meta states. 97mTc is the most stable, with a half-life of 90.1 days (0.097 MeV). This is followed by 95mTc (half life: 61 days, 0.038 MeV), and 99mTc (half-life: 6.01 hours, 0.143 MeV). 99mTc only emits gamma rays, subsequently decaying to 99 Tc.[26] For isotopes lighter than the most stable isotope, 98Tc, the primary decay mode is electron capture, giving molybdenum. For the heavier isotopes, the primary mode is beta emission, giving ruthenium, with the exception that 100Tc can decay both by beta emission and electron capture.[26] [27] Technetium-99 is the most common and most readily available isotope, as it is a major product of the fission of uranium-235. One gram of 99Tc produces 6.2×108 disintegrations a second (that is, 0.62 GBq/g).[4]
8
Technetium
9
Applications Nuclear medicine Technetium-99m or 99mTc ("m" indicates that this is a metastable nuclear isomer) is used in radioactive isotope medical tests, for example as a radioactive tracer that medical equipment can detect in the human body.[16] It is well suited to the role because it emits readily detectable 140 keV gamma rays, and its half-life is 6.01 hours (meaning that about fifteen sixteenths of it decays to 99Tc in 24 hours).[4] Klaus Schwochau's book Technetium lists 31 radiopharmaceuticals based on 99mTc for imaging and functional studies of the brain, myocardium, thyroid, lungs, liver, gallbladder, kidneys, skeleton, blood, and tumors.[28] Immunoscintigraphy incorporates
99m
Tc into a monoclonal
antibody, an immune system protein, capable of binding to cancer cells. A few hours after injection, medical equipment is used to detect the gamma rays emitted by the 99mTc; higher concentrations indicate where the tumor is. This technique is particularly useful for detecting hard-to-find cancers, such as those affecting the intestine. These modified antibodies are sold by the German company Hoechst (now part of Sanofi-Aventis) under the name "Scintium".[16] When
99m
Tc is combined with a tin compound it binds to
Technetium scintigraphy of a neck
red blood cells and can therefore be used to map of Graves' disease patient circulatory system disorders. It is commonly used to detect gastrointestinal bleeding sites. A pyrophosphate ion with 99m Tc adheres to calcium deposits in damaged heart muscle, making it useful to gauge damage after a heart attack.[29] The sulfur colloid of 99mTc is scavenged by the spleen, making it possible to image the structure of the spleen.[4] Radiation exposure due to diagnostic treatment involving Tc-99m can be kept low. Because 99m Tc has a short half-life and emits primarily a gamma ray (allowing small amounts to be easily detected), its quick decay into the far-less radioactive 99Tc results in relatively less total radiation dose to the patient, per unit of initial activity after administration. In the form administered in these medical tests (usually pertechnetate) both isotopes are quickly eliminated from the body, generally within a few days.[29] Technetium for nuclear medicine purposes is usually extracted from technetium-99m generators, because of its short 6 hour half-life.[30]
Technetium
Environmental science and biology The longer-lived isotope 95mTc, with a half-life of 61 days, is used as a radioactive tracer to study the movement of technetium in the environment and in plant and animal systems.[31]
Industrial and chemical Technetium-99 decays almost entirely by beta decay, emitting beta particles with consistent low energies and no accompanying gamma rays. Moreover, its long half-life means that this emission decreases very slowly with time. It can also be extracted to a high chemical and isotopic purity from radioactive waste. For these reasons, it is a NIST standard beta emitter, used for equipment calibration.[32] Technetium-99 has also been proposed for use in optoelectronic devices and nanoscale nuclear batteries.[33] Like rhenium and palladium, technetium can serve as a catalyst. For certain reactions, for example the dehydrogenation of isopropyl alcohol, it is a far more effective catalyst than either rhenium or palladium. Of course, its radioactivity is a major problem in finding safe applications.[34] Under certain circumstances, a small concentration (5×10−5 mol/L) of the pertechnetate ion in water can protect iron and carbon steels from corrosion. For this reason, pertechnetate has been used as a possible anodic corrosion inhibitor for steel, although technetium's radioactivity poses problems which limit this application to self-contained systems.[35] While (for example) CrO42− can also inhibit corrosion, it requires a concentration ten times as high. In one experiment, a test specimen was kept in an aqueous solution of pertechnetate for 20 years and was still uncorroded. The mechanism by which pertechnetate prevents corrosion is not well-understood, but seems to involve the reversible formation of a thin surface layer. One theory holds that the pertechnetate reacts with the steel surface to form a layer of technetium dioxide which prevents further corrosion; the same effect explains how iron powder can be used to remove pertechnetate from water. (Activated carbon can also be used for the same effect.) The effect disappears rapidly if the concentration of pertechnetate falls below the minimum concentration or if too high a concentration of other ions is added. As noted, the radioactive nature of technetium (3 MBq per liter at the concentrations required) makes this corrosion protection impractical in almost all situations. Nevertheless, corrosion protection by pertechnetate ions was proposed (but never adopted) for use in boiling water reactors.[36]
10
Technetium
11
Compounds and chemical reactions Chemical properties of technetium resemble those of rhenium, in particular chemical inertness and tendency to form covalent bonds. Contrary to manganese, technetium does not readily form cations.
Hydride Reaction of technetium with hydrogen produces anionic, negatively charged hydride [TcH9]2−. It consists of a trigonal prism with Tc atom in the center and six hydrogen atoms at the corners. Three more hydrogens make a triangle lying parallel to the base and crossing the prism in its center (see figure). A few hydrogens (~2) can be replaced by sodium (Na+) or potassium (K+) ions. [37]
Oxides Many technetium oxides are known. At temperatures 400-450 °C, technetium oxidizes to form pale-yellow heptoxide: 4 Tc + 7 O → 2 Tc O 2
It
adopts
a
2
Technetium hydride
7
centrosymmetric
corner-shared
bi-tetrahedral structure, in which the terminal and bridging Tc-O bonds are 167 pm and 184 pm respectively and the O-Tc-O angle is 180°[38] Technetium heptoxide is the anhydride of pertechnic acid and the precursor to sodium pertechnetate:[39] Tc2O7 + 2 NaOH → 2 NaTcO4 + H2O Black-colored technetium dioxide (TcO2) can be produced by reduction of heptoxide with technetium or hydrogen. [40]
Pertechnetic acid Pertechnetic acid (HTcO4) is produced by reacting Tc2O7 with water or oxidizing acids, such as nitric acid, concentrated sulfuric acid, aqua regia, or a mixture of nitric and hydrochloric acids. The resulting dark red, hygroscopic substance is a strong acid and easily donates protons. The remaining pertechnate anion TcO4− consists of a tetrahedron with oxygens in the corners and Te atom in the center. Unlike permanganate MnO4−, it is a weak oxidizing agent. Pertechnate is often used as a convenient water-soluble source of Te isotopes, such as Tc-99m, and as a catalyst. [41]
Technetium
12
Sulfide, selenide, telluride Technetium forms various sulfides. TcS2 is obtained by direct reaction of technetium and sulfur, and Tc2S7 is formed as follows: 2 HTcO4 + 7 H2S → Tc2S7 + 8 H2O In this reaction, technetium is not reduced, which is different from the similar reaction of manganese. Upon heating, technetium heptasulfide decomposes into disulphide and elementary sulfur: Tc2S7 → 2 TcS2 + 3 S Analogous reactions occur with selenium and tellurium.
Technetium clusters Tc6 and Tc8
[42]
Technetium clusters Two main technetium clusters are known, Tc6 and Tc8. Both clusters have prism shapes where vertical pairs of Tc atoms are connected by triple bonds and the planar atoms by single bonds (see figure) . Every Tc atoms makes six bonds the remaining valence electrons can be saturated by one axial and two mu-bridging halogen atoms such as chlorine or bromine. [43]
Organic complexes
Technetium carbonyl Tc2(CO)10
Technetium forms numerous organic complexes, which are relatively well investigated because of their importance for nuclear medicine. Technetium carbonyl Tc2(CO)10 is a white solid. In the molecule, two technetium atoms are weakly bound to each other; each atom is surrounded by octahedra of five carbonyl ligands (see left figure). The bond length between Tc atoms, 303 pm, is larger than the distance between two atoms in metallic technetium. Similar carbonyls are formed by manganese and rhenium. [44] An example of a technetium complex with an organic ligand is shown in the right figure and is used in nuclear medicine. It has a unique Tc-O moiety oriented normal to the plane of the molecule, where oxygen atom can be replaced by nitrogen.
Organic complex of technetium
Further reading • The radiochemical Manual, 2nd Ed, edited by B.J.
Wilson, 1966. • E.R.Scerri (2007). The Periodic Table, Its Story and Its Significance [45]. Oxford University Press,. http:/ / www. us. oup. com/ us/ catalog/ general/ subject/ Chemistry/ ?view=usa& ci=9780195305739. • Radiochemistry and nuclear chemistry [46], Gregory Choppin, Jan-Olov Liljenzin, and Jan Rydberg, 3rd Edition, 2002, the chapter on nuclear stability (pdf) [47] (viewed 5 January
Technetium 2007) • WebElements.com – Technetium [48], and EnvironmentalChemistry.com – Technetium [49] per the guidelines at Wikipedia's WikiProject Elements [50] (all viewed 1 December 2002) • Nudat 2 [51] nuclide chart from the National Nuclear Data Center, Brookhaven National Laboratory • Nuclides and Isotopes [52] Fourteenth Edition: Chart of the Nuclides, General Electric Company, 1989
References [1] " Technetium: technetium(III) iodide compound data (http:/ / openmopac. net/ data_normal/ technetium(iii) triiodide_jmol. html)". OpenMOPAC.net. . Retrieved 2007-12-10. [2] " Technetium: technetium(I) fluoride compound data (http:/ / www. openmopac. net/ data_normal/ tcfr_jmol. html)". OpenMOPAC.net. . Retrieved 2007-12-10. [3] C. R. Hammond. The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304857. [4] S. J. Rimshaw (1968). Cifford A. Hampel. ed. The Encyclopedia of the Chemical Elements. New York: Reinhold Book Corporation. pp. 689-693. [5] "Line Spectra of the Elements". The CRC Handbook. CRC press. 2004–2005. [6] *Schwochau, Klaus (2000). Technetium. Wiley-VCH. pp. 96. ISBN 3-527-29496-1. [7] S. H. Autler. " Technetium as a Material for AC Superconductivity Applications (http:/ / www. bnl. gov/ magnets/ Staff/ Gupta/ Summer1968/ 0049. pdf)". Proceedings of the 1968 Summer Study on Superconducting Devices and Accelerators. . Retrieved 2009-05-05. [8] Schwochau, pp. 371-381 [9] Schwochau, p.40 [10] Norman E. Holden. " History of the Origin of the Chemical Elements and Their Discoverers (http:/ / www. nndc. bnl. gov/ content/ elements. html)". Brookhaven National Laboratory. . Retrieved 2009-05-05. [11] Yoshihara, H. K. (2004). "Discovery of a new element 'nipponium': re-evaluation of pioneering works of Masataka Ogawa and his son Eijiro Ogawa". Atomic spectroscopy (Spectrochim. Acta, Part B) vol. 59 (no8): 1305–1310. doi: 10.1016/j.sab.2003.12.027 (http:/ / dx. doi. org/ 10. 1016/ j. sab. 2003. 12. 027). [12] Peter van der Krogt. " Elentymolgy and Elements Multidict, "Technetium" (http:/ / www. vanderkrogt. net/ elements/ elem/ tc. html)". . Retrieved 2009-05-05. [13] Armstrong, John T. (2003). " Technetium (http:/ / pubs. acs. org/ cen/ 80th/ technetium. html)". Chemical & Engineering News. . [14] Nies, Kevin A. (2001). " Ida Tacke and the warfare behind the discovery of fission (http:/ / www. hypatiamaze. org/ ida/ tacke. html)". . Retrieved 2009-05-05. [15] Using first-principles X-ray-emission spectral-generation algorithms developed at NIST, I simulated the X-ray spectra that would be expected for Van Assche's initial estimates of the Noddacks' residue compositions. The first results were surprisingly close to their published spectrum! Over the next couple of years, we refined our reconstruction of their analytical methods and performed more sophisticated simulations. The agreement between simulated and reported spectra improved further. Our calculation of the amount of element 43 required to produce their spectrum is quite similar to the direct measurements of natural technetium abundance in uranium ore published in 1999 by Dave Curtis and colleagues at Los Alamos. We can find no other plausible explanation for the Noddacks' data than that they did indeed detect fission "masurium.#Armstrong, John T. " Technetium" (http:/ / pubs. acs. org/ cen/ 80th/ technetium. html) Chemical & Engineering News (2003). [16] John Emsley (2001). Nature's Building Blocks: An A-Z Guide to the Elements. New York: Oxford University Press. ISBN 0-19-850340-7. [17] The transuranium people: The inside story (http:/ / www. worldscibooks. com/ physics/ p074. html). Chapter 1.2: Early Days at the Berkeley Radiation Laboratory: University of California, Berkeley & Lawrence Berkeley National Laboratory. 2000. pp. pp.15. ISBN ISBN 1-86094-087-0. . [18] Schwochau, pp.7-9 [19] Schwochau, pp.374-404 [20] K. Yoshihara (1996). "Technetium in the Environment". in K. Yoshihara and T. Omori. Topics in Current Chemistry: Technetium and Rhenium. 176. Berlin Heidelberg: Springer-Verlag. [21] Keiko Tagami (2003). " Technetium-99 Behaviour in the Terrestrial Environment - Field Observations and Radiotracer Experiments (http:/ / www. soc. nii. ac. jp/ jnrs/ paper/ JN41/ j041Tagami. pdf)". Journal of Nuclear
13
Technetium and Radiochemical Sciences 4: A1-A8. . [22] J D Harrison et al. (2001). " Gut transfer and doses from environmental technetium (http:/ / www. iop. org/ EJ/ abstract/ 0952-4746/ 21/ 1/ 004)". J. Radiol. Prot. 21: 9. doi: 10.1088/0952-4746/21/1/004 (http:/ / dx. doi. org/ 10. 1088/ 0952-4746/ 21/ 1/ 004). . [23] Arokiasamy J. Francis, Cleveland J. Dodge, G. E. Meinken. (2002). " Biotransformation of pertechnetate by Clostridia (http:/ / www. extenza-eps. com/ OLD/ doi/ abs/ 10. 1524/ ract. 2002. 90. 9-11_2002. 791)". Radiochimica Acta 90: 791. doi: 10.1524/ract.2002.90.9-11_2002.791 (http:/ / dx. doi. org/ 10. 1524/ ract. 2002. 90. 9-11_2002. 791). . [24] Schwochau, pp.87-96 [25] J. L. Snelgrove et al., (1995). " Development and Processing of LEU Targets for Mo-99 Production (http:/ / www. rertr. anl. gov/ MO99/ JLS. pdf)". . Retrieved 2009-05-05. [26] " "Technetium", Nuclides / Isotopes (http:/ / environmentalchemistry. com/ yogi/ periodic/ Tc-pg2. html)". . Retrieved 2009-05-05. [27] "Table of the isotopes". The CRC Handbook. CRC press. 2004–2005. [28] Schwochau, p.414 [29] Joseph F. Smith. " Technetium heart scan (http:/ / www. chclibrary. org/ micromed/ 00067370. html)". . Retrieved 2009-05-05. [30] Dilworth, Jonathan R.; Parrott, Suzanne J. (1998). "The biomedical chemistry of technetium and rhenium". Chemical Society Reviews 27: 43–55. doi: 10.1039/a827043z (http:/ / dx. doi. org/ 10. 1039/ a827043z). [31] Schwochau, pp. 12-27 [32] Schwochau, p. 87 [33] " University Research Program in Robotics REPORT (http:/ / www. osti. gov/ bridge/ servlets/ purl/ 895620-n4Nt3U/ 895620. PDF)" (PDF). University of Florida. 2006-11-30. . Retrieved 2007-10-12. [34] Schwochau, pp.87-90 [35] " EPA: 402-b-04-001b-14-final (http:/ / www. epa. gov/ radiation/ docs/ marlap/ 402-b-04-001b-14-final. pdf)" (PDF). Marlap. July 2004. . Retrieved 2008-08-04. [36] Schwochau, p.91 [37] Schwochau, p.104 [38] Krebs, B. (1969). "Technetium(VII)-oxid: Ein Übergangsmetalloxid mit Molekülstruktur im festen Zustand". Angewandte Chemie 81: 328 - 329. doi: doi (http:/ / dx. doi. org/ doi). [39] Herrell, A. Y.; Busey, R. H.; Gayer, K. H. (1977). Technetium(VII) Oxide, in Inorganic Syntheses. XVII. pp. 155-158. ISBN 0-07-044327-0. [40] Schwochau, p.108 [41] Schwochau, pp.127-136 [42] Schwochau, pp.112-113 [43] German, K. E.; Kryutchkov S. V. (2002). "Polynuclear Technetium Halide Clusters". Russian Journal of Inorganic Chemistry 47 (4): 578–583.. [44] [45] [46] [47] [48] [49] [50] [51] [52]
Schwochau, p.286, 328 http:/ / www. us. oup. com/ us/ catalog/ general/ subject/ Chemistry/ ?view=usa& ci=9780195305739 http:/ / book. nc. chalmers. se/ http:/ / book. nc. chalmers. se/ KAPITEL/ CH03NY3. PDF http:/ / www. webelements. com/ webelements/ elements/ text/ Tc/ index. html http:/ / environmentalchemistry. com/ yogi/ periodic/ Tc. html http:/ / en. wikipedia. org/ wiki/ Wikipedia:WikiProject_Elements http:/ / www. nndc. bnl. gov/ nudat2/ index. jsp http:/ / chartofthenuclides. com/ default. html
14
Article Sources and Contributors
Article Sources and Contributors Technetium Source: http://en.wikipedia.org/w/index.php?oldid=307534660 Contributors: -- April, 2over0, Aarchiba, Adamrush, Aditya, Ahoerstemeier, AjAldous, Alexius08, Alfio, Alhizar, Ali Graham, AlimanRuna, Andres, Anthony Appleyard, Anville, Apostrophe, Archimerged, Arkrishna, ArnoldReinhold, Arvinkong, Astavats, AstroHurricane001, Avg, Bantman, Barticus88, Bazzargh, Bcrowell, Beetstra, Benbest, Binary TSO, Bk0, Blahblahblahblah01, BlankVerse, BlueEarth, Bm gub, BorgQueen, Brighterorange, Bryan Derksen, CTZMSC3, CYD, Cadmium, Calliopejen1, Carnildo, Cedrus-Libani, CesarB, Chanting Fox, Chrislk02, Cmdrjameson, Coemgenus, Conscious, Conversion script, Cryptic C62, CryptoDerk, Cstaffa, DMacks, DV8 2XL, Dale Arnett, Darrien, David Latapie, Dcoetzee, Deglr6328, Dekisugi, Delta G, DerHexer, Dirac66, Disavian, Dmn, Donarreiskoffer, Doodledoo, DragonflySixtyseven, Drchessman, Dreish, Dwmyers, Edgar181, El C, Eleassar777, Eleuther, Emperorbma, Eog1916, Everyking, Evil Monkey, Eweisser, Femto, Fibonacci, Firien, Flypro, Fredrik, Fresheneesz, Gaius Cornelius, GamblinMonkey, Gene Nygaard, Geni, Giftlite, Gmaxwell, Gravitan, GreatMizuti, Greybeard, Gseryakov, H Padleckas, Hadal, Hak-kâ-ngìn, Hashproduct, Hede2000, Hellbus, HereToHelp, Ian Pitchford, Icairns, Ideyal, Imperator Honorius, InternetMeme, J.delanoy, JRM, JWB, JYolkowski, Jan van Male, Jaraalbe, Jerzy, Jimp, Jkl, Joanjoc, Joe Shupienis, John, John254, Jose77, Joshschr, Joshuaali, Judgesurreal777, Kaal, Karl-Henner, Keenan Pepper, Khaosworks, Kira134341, KnowledgeOfSelf, Konstantin.German, Kurykh, Kwamikagami, La goutte de pluie, Lament, LarryMorseDCOhio, LeaveSleaves, Lightmouse, Ligulem, Luigi30, Lupin, Malcolm Farmer, Materialscientist, Matt Gies, Mattd4u2nv, Mav, Mejor Los Indios, Mgimpel, MickWest, Mike Clough, Minesweeper, Mjpieters, Morbid-o, Mr.Z-man, Neckro, Nedffred, Nergaal, Nihiltres, Nikai, Nilmerg, Oblivious, Oxymoron83, Pacotaco321, Pakaran, Peruvianllama, Pete Julishman, Petri Krohn, PhilKnight, Physchim62, Piano non troppo, Pixel ;-), PlatinumX, Plexust, Poolkris, Q43, Quadell, RAM, Ray Van De Walker, Rdsmith4, RedHillian, Reid, Reinyday, Remember, Res2216firestar, Reyk, Rich Farmbrough, Rintrah, Rjreliga, Rjwilmsi, RobertG, Roberta F., Romanm, Rursus, RxS, SWAdair, Sahands, Saperaud, Sat84, Sbharris, Schneelocke, Sengkang, Sfahey, Shafei, Shanes, Shimgray, Sjakkalle, Skizzik, Sl, Snagglepuss, Snydley, Spiff, Squids and Chips, StaticGull, Stemonitis, Stifynsemons, Stone, Suisui, Sunborn, Sunnyoraish, Svante, TUF-KAT, Tagishsimon, Taxman, Tetracube, Trovatore, Ttony21, Ttwaring, Until It Sleeps, VASANTH S.N., Vsmith, Warut, Watch37264, Wayward, WindOwl, Xezbeth, Xmnemonic, Yvwv, Yyy, Zephyr2k, Zfr, Zoicon5, 312 anonymous edits
Image Sources, Licenses and Contributors image:Tc-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Tc-TableImage.png License: GNU Free Documentation License Contributors: Daniel Mayer and Arnaud Gaillard Image: Tc,43.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Tc,43.jpg License: Public Domain Contributors: Kilom691, Ms2ger, Saperaud, 1 anonymous edits Image:Дмитрий Иванович Менделеев 4.gif Source: http://en.wikipedia.org/w/index.php?title=File:Дмитрий_Иванович_Менделеев_4.gif License: Public Domain Contributors: Anrie, Maximaximax, OldakQuill Image:UraniumUSGOV.jpg Source: http://en.wikipedia.org/w/index.php?title=File:UraniumUSGOV.jpg License: Public Domain Contributors: Kluka, Saperaud Image:Basedow-vor-nach-RIT.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Basedow-vor-nach-RIT.jpg License: Copyrighted free use Contributors: Al.locke, Nard the Bard, Saperaud, 1 anonymous edits Image:Technetiumhydrid.png Source: http://en.wikipedia.org/w/index.php?title=File:Technetiumhydrid.png License: Public Domain Contributors: Aglarech Image:Technetiumcluster.png Source: http://en.wikipedia.org/w/index.php?title=File:Technetiumcluster.png License: Public Domain Contributors: Aglarech, Mdd Image:Technetiumcarbonyl.png Source: http://en.wikipedia.org/w/index.php?title=File:Technetiumcarbonyl.png License: Public Domain Contributors: Aglarech, Benjah-bmm27 Image:Technetiumkomplex.png Source: http://en.wikipedia.org/w/index.php?title=File:Technetiumkomplex.png License: Public Domain Contributors: Aglarech, Benjah-bmm27
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15
Ruthenium
1
Ruthenium 44
technetium ← ruthenium → rhodium
Fe ↑
Ru ↓
Os Periodic Table - Extended Periodic Table
General Name, Symbol, Number
ruthenium, Ru, 44
Element category
transition metals
Group, Period, Block
8, 5, d
Appearance
silvery white metallic
Standard atomic weight
101.07(2) g·mol
Electron configuration
[Kr] 4d 5s
−1
7
Electrons per shell
1
2, 8, 18, 15, 1 Physical properties
Density (near r.t.)
12.45 g·cm
Liquid density at m.p.
10.65 g·cm−3
Melting point
2607 K (2334 °C, 4233 °F)
Boiling point
4423 K (4150 °C, 7502 °F)
Heat of fusion
38.59 kJ·mol−1
Heat of vaporization
591.6 kJ·mol−1
Specific heat capacity
(25 °C) 24.06 J·mol−1·K−1
−3
Vapor pressure P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
2588
2811
3087
3424
3845
4388
Atomic properties Crystal structure
hexagonal
Ruthenium
Oxidation states
Electronegativity Ionization energies
2 [1]
8, 7, 6, 4, 3, 2, 1, , -2 (mildly acidic oxide) 2.3 (Pauling scale) 1st: 710.2 kJ/mol 2nd: 1620 kJ/mol 3rd: 2747 kJ/mol
Atomic radius
134 pm
Covalent radius
146±7 pm Miscellaneous
Magnetic ordering Electrical resistivity
[2]
paramagnetic
(0 °C) 71 nΩ·m
Thermal conductivity
(300 K) 117 W·m
Thermal expansion
(25 °C) 6.4 µm·m
Speed of sound (thin rod)
(20 °C) 5970 m/s
Young's modulus
447 GPa
Shear modulus
173 GPa
Bulk modulus
220 GPa
−1
−1
Poisson ratio
0.30
Mohs hardness
6.5
Brinell hardness CAS registry number
·K
2160 MPa 7440-18-8 Most-stable isotopes
−1
·K
−1
Ruthenium
3
Main article: Isotopes of ruthenium iso 96
Ru
97
Ru
NA
half-life
5.52% syn
2.9 d
98
99
12.7%
99
100
12.6%
100
101
17.0%
101
102
31.6%
102
Ru Ru Ru
103
Ru
104
Ru
106
Ru
DP
Ru is stable with 52 neutron
1.88%
Ru
DE (MeV)
96
98
Ru
DM
ε
-
γ
0.215, 0.324
97
Tc
-
Ru is stable with 54 neutron Ru is stable with 55 neutron Ru is stable with 56 neutron Ru is stable with 57 neutron Ru is stable with 58 neutron
syn
39.26 d
18.7%
β
0.226
γ
0.497
−
103
Rh
-
104
Ru is stable with 60 neutron
syn
373.59 d
3.54
−
β
106
Rh
References
Ruthenium (pronounced /ruːˈθiːniəm/) is a chemical element that has the symbol Ru and atomic number 44. A rare transition metal of the platinum group of the periodic table, ruthenium is found associated with platinum ores and used as a catalyst in some platinum alloys.
Characteristics A polyvalent hard white metal, ruthenium is a member of the platinum group, and is in group 8 of the periodic table: Z
Element
No. of electrons/shell
26
iron
2, 8, 14, 2
44
ruthenium
2, 8, 18, 15, 1
76
osmium
2, 8, 18, 32, 14, 2
108
hassium
2, 8, 18, 32, 32, 14, 2
but has an atypical configuration in its outermost electron shells compared to the rest of the members. (This can be observed in the neighborhood of niobium (41), ruthenium (44), rhodium (45), and palladium (46).) Ruthenium has four crystal modifications and does not tarnish at normal temperatures, but does oxidize readily on exposure to air to form ruthenium tetroxide, RuO4, a strong oxidising agent with properties analogous to those of osmium tetroxide. Ruthenium dissolves in fused alkalis, is not attacked by acids but is attacked by halogens at high temperatures. Small amounts of ruthenium can increase the hardness of platinum and
Ruthenium palladium. The corrosion resistance of titanium is increased markedly by the addition of a small amount of ruthenium. This metal can be plated either by electroplating or by thermal decomposition methods. One ruthenium-molybdenum alloy has been found to be superconductive at 10.6 K. The oxidation states of ruthenium range from +1 to +8, and -2 is known, though oxidation states of +2, +3, and +4 are most common.
Uses Due to its ability to harden platinum and palladium, ruthenium is used in platinum and palladium alloys to make wear-resistant electrical contacts. Because of its lower cost and similar properties compared to rhodium,[3] the use as plating material for electric contacts is one of the major applications.[4] [5] The coatings are either put on by electroplating[6] or sputtering[7] . Ruthenium dioxide and lead and bismuth[8] ruthenates, the later with perovskite crystal structure,[9] are used thick film chip resistors. A mixture of the one of the above mentioned substances is mixed with a glass substrate and printed on a ceramic material and heated to 860°C.[10] The first two applications account for 50% of the ruthenium consumption.[11] It is sometimes alloyed with gold in jewelry. 0.1% ruthenium is added to titanium to improve its corrosion resistance.[12] Ruthenium is also used in some advanced high-temperature single-crystal superalloys, with applications including the turbine blades in jet engines. Two nickel based superalloy compositions described in the literature are EPM-102 (with 3 % Ru) and TMS-162 (with 6 % Ru), both contain 6 % rhenium.[13] Fountain pen nibs are frequently tipped with alloys containing ruthenium. From 1944 onward, the famous Parker 51 fountain pen was fitted with the "RU" nib, a 14K gold nib tipped with 96.2% ruthenium and 3.8% iridium. Ruthenium is also a versatile catalyst. Hydrogen sulfide can be split by light by using an aqueous suspension of CdS particles loaded with ruthenium dioxide. This may be useful in the removal of H2S from oil refineries and from other industrial processes. Ruthenium is a component of mixed-metal oxide (MMO) anodes used for cathodic protection of underground and submerged structures, and for electrolytic cells for chemical processes such as generating chlorine from saltwater. Organometallic ruthenium carbene and allenylidene complexes have recently been found as highly efficient catalysts for olefin metathesis with important applications in organic and pharmaceutical chemistry. The fluorescence of some ruthenium complexes is quenched by oxygen, which has led to their use as optode sensors for oxygen. Ruthenium red, [(NH3)5Ru-O-Ru(NH3)4-O-Ru(NH3)5]6+, is a biological stain used to stain polyanionic molecules such as pectin and nucleic acids for light microscopy and electron microscopy. The beta-decaying isotope 106 of ruthenium is used in radiotherapy of eye tumors, mainly malignant melanomae of the uvea. Ruthenium-centered complexes are being researched for possible anticancer properties.[14] Ruthenium, unlike traditional platinum complexes, show greater resistance to hydrolysis
4
Ruthenium and more selective action on tumors. NAMI-A and KP1019 are two drugs undergoing clinical evaluation against metastatic tumors and colon cancers.
Applications of ruthenium thin films in microelectronics Relatively recently, ruthenium has been suggested as a material that could beneficially replace other metals and silicides in microelectronics components. Ruthenium tetroxide (RuO4) is highly volatile, as is ruthenium trioxide (RuO3).[15] By oxidizing ruthenium (for example with an oxygen plasma) into the volatile oxides, ruthenium can be easily patterned.[16] [17] [18] [19] The properties of the common ruthenium oxides make ruthenium a metal compatible with the semiconductor processing techniques needed to manufacture microelectronics. In order to continue miniaturization of microelectronics, new materials are needed as dimensions change. There are three main applications for thin ruthenium films in microelectronics. The first is using thin films of ruthenium as electrodes on both sides of tantalum pentoxide (Ta2O5) or barium strontium titanate ((Ba, Sr)TiO3, also known as BST) in the next generation of three-dimensional dynamic random access memories (DRAMs).[20] [21] [22] Ruthenium thin film electrodes could also be deposited on top of lead zirconate titanate (Pb(ZrxTi1-x)O3, also known as PZT) in another kind of RAM, ferroelectric random access memory (FRAM).[23] [24] Platinum has been used as the electrodes in RAMs in laboratory settings, but it is difficult to pattern. Ruthenium is chemically similar to platinum, preserving the function of the RAMs, but in contrast to Pt patterns easily. The second is using thin ruthenium films as metal gates in p-doped metal-oxide-semiconductor field effect transistors (p-MOSFETs).[25] When replacing silicide gates with metal gates in MOSFETs, a key property of the metal is its work function. The work function needs to match the surrounding materials. For p-MOSFETs, the ruthenium work function is the best materials property match with surrounding materials such as HfO2, HfSiOx, HfNOx, and HfSiNOx, to achieve the desired electrical properties. The third large-scale application for ruthenium films is as a combination adhesion promoter and electroplating seed layer between TaN and Cu in the copper dual damascene process.[26] [27] [28] [29] [30] Copper can be directly electroplated onto ruthenium[31] , in contrast to tantalum nitride. Copper also adheres poorly to TaN, but well to Ru. By depositing a layer of ruthenium on the TaN barrier layer, copper adhesion would be improved and deposition of a copper seed layer would not be necessary. There are also other suggested uses. In 1990, IBM scientists discovered that a thin layer of ruthenium atoms created a strong anti-parallel coupling between adjacent ferromagnetic layers, stronger than any other nonmagnetic spacer-layer element. Such a ruthenium layer was used in the first giant magnetoresistive read element for hard disk drives. In 2001, IBM announced a three-atom-thick layer of the element ruthenium, informally referred to as "pixie dust", which would allow a quadrupling of the data density of current hard disk drive media.[32]
5
Ruthenium
Thin-film Solar cells Some ruthenium complexes absorb light throughout the visible spectrum and are being actively researched in various, potential, solar energy technologies. Ruthenium-based dyes have been used as the electron providers in dye-sensitized solar cells, a promising new low-cost solar cell system.
History Though naturally occurring platinum, containing all six platinum group metals, was used for a long time by pre-Columbian Americans and known to European chemists from the mid-16th century, it took until the mid-17th century for platinum to be discovered. The discovery that natural platinum contained palladium, rhodium, osmium and iridium took place in the first decade of the 18th century. Platinum in alluvial sands of Russian rivers gave access to raw material for use in plates and medals and for the minting of ruble coins, starting in 1828.[33] Residues of platinum production for minting where available in the Russian Empire, and therefore most of the research on them was done in Eastern Europe. It is possible that the Polish chemist Jędrzej Śniadecki isolated element 44 (which he called "vestium") from platinum ores in 1807. His work was never confirmed, however, and he later withdrew his claim of discovery. Jöns Berzelius and Gottfried Osann nearly discovered ruthenium in 1827.[34] The men examined residues that were left after dissolving crude platinum from the Ural Mountains in aqua regia. Berzelius did not find any unusual metals, but Osann thought he found three new metals, pluranium, ruthenium and polinium. This discrepency led to a long-standing controversy between Berzelius and Osann about the composition of the residues.[35] In 1844 the Russian scientist Karl Klaus showed that the compounds prepared by Gottfried Osann contained small amounts of ruthenium, which Klaus had discovered the same year.[36] Klaus isolated ruthenium from the platinium residues of the rouble production while he was working in Kazan University, Kazan.[35] Klaus showed that ruthenium oxide contained a new metal and obtained 6 grams of ruthenium from the part of crude platinum that is insoluble in aqua regia.[35] The name derives from Ruthenia, the Latin word for Rus', a historical area which includes present-day western Russia, Ukraine, Belarus, and parts of Slovakia and Poland. Karl Klaus named the element in honour of his birthland, as he was born in Tartu, Estonia, which was at the time a part of the Russian Empire.
Occurrence Normal mining This element is generally found in ores with the other platinum group metals in the Ural Mountains and in North and South America. Small but commercially important quantities are also found in pentlandite extracted from Sudbury, Ontario, Canada, and in pyroxenite deposits in South Africa. The native ruthenium is very rare mineral (Ir replaces part of Ru in its structure).[4] [37] Ruthenium is exceedingly rare and is the 74th most abundant metal on Earth.[11] Roughly 12 tonnes of Ru is mined each year with world reserves estimated to be 5000 tonnes.[11]
6
Ruthenium The composition of the mined platinum group metal mixtures varies in a wide range depending on the geochemical formation, for example the PGM mined in South Africa contain on average 11% ruthenium while the PGM mined in the USSR containe only 2% based on reseach dating from 1992.[38] [39]
From used nuclear fuels It is also possible to extract ruthenium from used nuclear fuel. Each kilo of fission products of 235U will contain 63.44 grams of ruthenium isotopes with halflives longer than a day. Since a typical used nuclear fuel contains about 3% fission products, one ton of used fuel will contain about 1.9 kg of ruthenium. The 103Ru and 106Ru will render the fission ruthenium very radioactive. If the fission occurs in an instant then the ruthenium thus formed will have an activity due to 103Ru of 109 TBq g-1 and 106Ru of 1.52 TBq g-1. Ru 103 has a half life of about 39 days meaning that within 390 days it will have effectively decayed to ground state, well before any reprocessing is likely to occur. Ru 106 has a half life of about 373 days meaning that if the fuel is let to cool for 5 years before reprocessing only about 3% of the original quantity will remain, the rest will have decayed to ground state.
Production Ruthenium, like the other platin group metals, is obtained commercially as a by-product from nickel and copper mining and processing or by direct processing of platin group metal ores. During electrorefining of copper and nickel, noble metals such as silver, gold and the platinum group metals The radioactivity in MBq per gram of each of the platinum group including selenium and metals which are formed by the fission of uranium. Of the metals tellurium settle to the bottom shown, ruthenium is the most radioactive. Palladium has an almost of the cell as anode mud, constant activity due to the very long lived 107Pd while rhodium is the least radioactive. which forms the starting point [37] [4] for their extraction. In order to separate the metals, they must first be brought into solution. Several methods are available depending on the separation process and the composition of the mixture; two representative methods are fusion with sodium peroxide followed by dissolution in aqua regia, and dissolution in a mixture of chlorine with hydrochloric acid.[40] [41] Osmium, ruthenium, rhodium and iridium can be separated from platinum and gold and base metals by their insolubility in aqua regia, leaving a solid residue. Rhodium can be separated from the residue by treatment with molten sodium bisulphate. The insoluble residue, containing Ru, Os and Ir is treated with sodium oxide, in which Ir is insoluble, producing water-soluble Ru and Os salts. After oxidation to the volatile oxides, RuO4 is separated from OsO4 by precipitation of (NH4)3RuCl6 with ammonium chloride.
7
Ruthenium This metal is commercially isolated through a complex chemical process in which hydrogen is used to reduce ammonium ruthenium chloride yielding a powder. The powder is then consolidated by powder metallurgy techniques or by argon-arc welding. After it is dissolved, osmium is separated from the other platinum group metals by distillation or extraction with organic solvents of the volatile osmium tetroxide.[42] The first method is similar to the procedure Tennant and William Hyde Wollaston used for their separation. Both methods are suitable for industrial scale production. In either case, the product is reduced using hydrogen, yielding the metal as a powder or sponge that can be treated using powder metallurgy techniques.[43]
Compounds Ruthenium compounds are often similar in properties to those of osmium and exhibit at least eight oxidation states, but the +2, +3, and +4 states are the most common. Examples are ruthenium(IV) oxide (RuO2, oxidation state +4), dipotassium ruthenate (K2RuO4, +6), potassium perruthenate (KRuO4, +7) and ruthenium tetroxide (RuO4, +8). Compounds of ruthenium with chlorine are ruthenium(II) chloride (RuCl2) and ruthenium(III) chloride (RuCl3).
Isotopes Naturally occurring ruthenium is composed of seven stable isotopes. Additionally, 34 radioactive isotopes have been discovered. Of these radioisotopes, the most stable are 106 Ru with a half-life of 373.59 days, 103Ru with a half-life of 39.26 days and 97Ru with a half-life of 2.9 days. Fifteen other radioisotopes have been characterized with atomic weights ranging from 89.93 u (90Ru) to 114.928 u (115Ru). Most of these have half-lives that are less than five minutes except 95Ru (half-life: 1.643 hours) and 105Ru (half-life: 4.44 hours). The primary decay mode before the most abundant isotope, 102Ru, is electron capture and the primary mode after is beta emission. The primary decay product before 102Ru is technetium and the primary mode after is rhodium.
Organometallic chemistry Ruthenium is a versatile metal that can easily form compounds with carbon ruthenium bonds, as these compounds tend to be darker and react more quickly than the osmium compounds. Recently, Professor Anthony Hill and his co-workers have been making compounds of ruthenium in which a boron atom binds to the metal atom.[44] The organometallic ruthenium compound that is easiest to make is RuHCl(CO)(PPh3)3. This compound has two forms (yellow and pink) that are identical once they are dissolved but different in the solid state. An organometallic compound similar to ruthenocene, bis(2,4-dimethylpentadienyl)ruthenium, is readily synthesized in near quantitative yields and has applications in vapor-phase deposition of metallic ruthenium, as well as in catalysis, including Fischer-Tropsch synthesis of transportation fuels. Important catalysts based on ruthenium are Grubbs' catalyst and Roper's complex.
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Ruthenium
Chemical vapor deposition of ruthenium A unique challenge arises in trying to grow impurity-free films of a catalyst in Chemical vapor deposition (CVD). Ruthenium metal activates C-H and C-C bonds, which aids C-H and C-C bond scission. This creates a potential catalytic decomposition path for all metal-organic CVD precursors that is likely to lead to significant carbon incorporation. Platinum, a chemically similar catalyst, catalyzes dehydrogenation of five- and six-member cyclic hydrocarbons into benzene.[45] The d-bands of ruthenium lie higher than those in platinum, generally predicting stronger ruthenium-adsorbate bonds than on platinum. Therefore, it is likely that ruthenium also catalyzes dehydrogenation of five- and six-member hydrocarbon rings to benzene. Benzene dehydrogenates further on ruthenium surfaces into hydrocarbon fragments similar to those formed by acetylene and ethene on ruthenium surfaces.[46] [47] In addition to benzene, acetylene and ethene, pyridine also decomposes on ruthenium surfaces, leaving bound fragments on the surface.[48] Ruthenium is unusually well studied in the surface science and catalysis literature due to its industrial importance as a catalyst. There are many studies of individual molecular behavior on ruthenium in surface science. However, understanding the behavior of each ligand on its own is not equivalent to understanding their behavior when co-adsorbed with each other and with the precursor. While there is no significant pressure difference between surface science studies and CVD, there is often a temperature gap between temperatures reported in surface science studies and CVD growth temperatures. Despite these complications, ruthenium is a promising candidate for understanding chemical vapor deposition and precursor design of catalytic films. Ligands that are stable compounds in their own right, short ligand-ruthenium contact times and moderate substrate temperatures help minimize unwanted ligand decomposition on the surface.[49] [50] [51] The C-H and C-C bond activation is temperature-dependent. Product desorption is also temperature-dependent, if the products are not bound to the ruthenium surface. This suggests that there is some optimum temperature, at which most independently stable ligands have just enough thermal energy to desorb from the ruthenium film surface before C-H activation can occur. For example, benzene starts decomposing on ruthenium at 87°C. However, the dehydrogenation reaction does not go to fragments until 277°C, and compete fragmentation is not seen at low surface coverages. This suggests that provided adsorbed benzene molecules are not close to one another on the surface and temperatures are below 277°C, the vast majority of benzene molecules may not contribute to carbon incorporation in films. Therefore, a key consideration in growing CVD films of catalytic metals such as ruthenium is combining molecule design and the kinetic aspects of growth in a favorable way. Before metal-organic precursors were explored, triruthenium dodecacarbonyl (Ru3(CO)12) was tested as a CVD precursor.[52] [53] While this precursor gives good-quality films, the vapor pressure is poor, complicating its practical use in a CVD process. Ruthenocene[54] [55] and bis(ethylcyclopentadienyl)ruthenium(II)[56] [57] [58] [59] and beta-diketonate ruthenium(II) compounds[60] [61] [62] have been fairly extensively explored. Although these precursors also can give pure films of low resistivity when reacted with oxygen, the growth rates are very low or not reported. One high-growth precursor, cyclopentadienyl-propylcyclopentadienylruthenium(II) (RuCp(i-PrCp)), has been identified.[63] (RuCp(i-PrCp) has achieved growth rates of 7.5 nm/min to 20 nm/min as well as low resistivities. However, it does not nucleate on oxides, ruling out its use in all
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Ruthenium applications but copper interconnect playing layers. A new zero-valent, single-source precursor design paradigm was launched by Schneider et al. with (1,5-cyclooctadiene)(toluene)Ru(0) ((1,5-COD)(toluene)Ru)[64] and [65] (1,3-cyclohexadiene)(benzene)Ru(0) ((1,3-CHD)(benzene)Ru) , also independently tested [66] by Choi et al. Using (1,5-COD)(toluene)Ru, Schneider found that C-H bonds were readily activated in 1,5-COD. Although carbon incorporation levels were low (1-3%), the growth rates were only around 0.28 nm/min at best. Using (1,3-CHD)(benzene)Ru, the 1,3-CHD was dehydrogenated to benzene as expected, but the large variety of possible surface reactions involving the two ligands resulted in a narrow process window in which carbon concentrations were low.
Precautions The compound ruthenium tetroxide, RuO4, similar to osmium tetroxide, is volatile, highly toxic and may cause explosions if allowed to come into contact with combustible materials.[67] Ruthenium plays no biological role but does strongly stain human skin, may be carcinogenic[68] and bio-accumulates in bone.
In fiction Ruthenium's use as a catalyst is a plot device in the novel Arctic Drift by Clive Cussler and Dirk Cussler
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A.; Dermann, K.; Rothaut, J.; Drieselman, R. (2002). "Platinum group metals and compounds". Ullmann's Encyclopedia of Industrial Chemistry. Wiley. doi: 10.1002/14356007.a21_075 (http:/ / dx. doi. org/ 10. 1002/ 14356007. a21_075). [41] Seymour, R. J.; O'Farrelly, J. I. (2001). "Platinum-group metals". Kirk Othmer Encyclopedia of Chemical Technology. Wiley. doi: 10.1002/0471238961.1612012019052513.a01.pub2 (http:/ / dx. doi. org/ 10. 1002/ 0471238961. 1612012019052513. a01. pub2). [42] Gilchrist, Raleigh (1943). "The Platinum Metals.". Chemical Reviews 32 (3): 277–372. doi: 10.1021/cr60103a002 (http:/ / dx. doi. org/ 10. 1021/ cr60103a002). [43] Hunt, L. B.; Lever, F. M. (1969). " Platinum Metals: A Survey of Productive Resources to industrial Uses (http:/ / www. platinummetalsreview. com/ pdf/ pmr-v13-i4-126-138. pdf)". Platinum Metals Review 13 (4): 126–138. . [44] - Professor Anthony Hill (http:/ / rsc. anu. edu. au/ research/ hill. php) - Current Research [45] Manner, W. L.; Girolami, G. S.; Nuzzo, R. G., Sequential Dehydrogenation of Unsaturated Cyclic C5 and C6 Hydrocarbons on Pt(111). J. Phys. Chem. B 1998, 102, (50), 10295-10306. [46] Jakob, P.; Menzel, D., The adsorption of benzene on Ru(001). Surface Science 1988, 210, 503-530. [47] Jakob, P. Adsorption und thermische Evolution von aromatischen Molekülen auf Ru(001): Benzol, Benzol+O, CO, H und Pyridin. Doctoral, Technische Universität München, München, 1989. [48] Jakob, P. Adsorption und thermische Evolution von aromatischen Molekülen auf Ru(001): Benzol, Benzol+O, CO, H und Pyridin. Doctoral, Technische Universität München, München, 1989. [49] Schneider, A.; Popovska, N.; Jipa, I.; Atakan, B.; Siddiqi, M. A.; Siddiqui, R.; Zenneck, U., Minimizing the carbon content of thin ruthenium films by MOCVD precursor complex design and process control. Chemical Vapor Deposition 2007, 13, (8), 389-395. [50] Schneider, A.; Popovska, N.; Holzmann, F.; Gerhard, H.; Topf, C.; Zenneck, U., [(1,5-Cyclooctadiene)(toluene)ruthenium(0)]: A Novel Precursor for the MOCVD of Thin Ruthenium Films. Chemical Vapor Deposition 2005, 11, (2), 99-105. [51] Schneider, A. Metallorganische Chemische Dasphasenabscheidung (MOCVD) von Übergangsmetallen am Beispiel von Eisen, Ruthenium und Wolfram. Doctoral, Universität Erlangen-Nürnberg, Erlangen, 2006. [52] Green, M. L.; Gross, M. L.; Papa, L. E.; Schnoes, K. J.; Brasen, D., Chemical Vapor Deposition of Ruthenium and Ruthenium Dioxide Films. Journal of The Electrochemical Society 1985, (132), 2677. [53] Wang, Q.; Ekerdt, J. G.; Gay, D.; Sun, Y.-M.; White, J. M., Low-temperature chemical vapor deposition and scaling limit of ultrathin Ru films. Applied Physics Letters 2004, 84, (8), 1380-1382. [54] Trent, D. E.; Paris, B.; Krause, H. H., Vapor Deposition of Pure Ruthenium Metal from Ruthenocene. Inorg. Chem. 1964, 3, (7), 1057-1058. [55] Park, S. E.; Kim, H. M.; Kim, K. B.; Min, S. H., Metallorganic chemical vapor deposition of Ru and RuO2 using ruthenocene precursor and oxygen gas. Journal of the Electrochemical Society 2000, 147, (1), 203-209. [56] Aoyama, T.; Eguchi, K., Ruthenium films prepared by liquid source chemical vapor deposition using bis-(ethylcyclopentadienyl)ruthenium. Japanese Journal of Applied Physics 1999, 38, (10A), 1134-6. [57] Kang, S. Y.; Choi, K. H.; Lee, S. K.; Hwang, C. S.; Kim, H. J., Thermodynamic Calculations and Metallorganic Chemical Vapor Deposition of Ruthenium Thin Films Using Bis(ethyl-pi-cyclopentadienyl)Ru for Memory Applications. Journal of The Electrochemical Society 2000, 147, (3), 1161-1167. [58] Matsui, Y.; Hiratani, M.; Nabatame, T.; Shimamoto, Y.; Kimura, S., Characteristics of Ruthenium Films Prepared by Chemical Vapor Deposition Using Bis(ethylcyclopentadienyl)ruthenium Precursor. Electrochemical and Solid-State Letters 2002, 5, (1), C18-C21. [59] Nabatame, T.; Hiratani, M.; Kadoshima, M.; Shimamoto, Y.; Matsui, Y.; Ohji, Y.; Asano, I.; Fujiwara, T.; Suzuki, T., Properties of ruthenium films prepared by liquid source metalorganic chemical vapor deposition using Ru(EtCp)2 with tetrahydrofuran solvent. Japanese Journal of Applied Physics 2000, 39, (11B), 1188-90. [60] Kadoshima, M.; Nabatame, T.; Hiratani, M.; Nakamura, Y.; Asano, I.; Suzuki, T., Ruthenium Films Prepared by Liquid Source Metalorganic Chemical Vapor Deposition Using Ru(dpm)3 Dissolved with Tetrahydrofuran Solvent. Japanese Journal of Applied Physics 2002, 41Part 2, (3B), L347-L350.
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Ruthenium [61] Lai, Y.-H.; Chen, Y.-L.; Chi, Y.; Liu, C.-S.; Carty, A. J.; Peng, S.-M.; Lee, G.-H., Deposition of Ru and RuO2 thin films employing dicarbonyl bis-diketonate ruthenium complexes as CVD source reagents. Journal of Materials Chemistry 2003, 13, 1999-2006. [62] Lee, J.-H.; Kim, J.-Y.; Rhee, S.-W.; Yang, D.; Kim, D.-H.; Yang, C.-H.; Han, Y.-K.; Hwang, C.-J., Chemical vapor deposition of Ru thin films by direct liquid injection of Ru(OD)3 (OD=octanedionate). Journal of Vacuum Science & Technology A 2000, 18, (5), 2400-2403. [63] Kang, S. Y.; Lim, H. J.; Hwang, C. S.; Kim, H. J., Metallorganic chemical vapor deposition of Ru films using cyclopentadienyl-propylcyclopentadienylruthenium(II) and oxygen. Journal of the Electrochemical Society 2002, 149, (6), C317-C323. [64] Schneider, A.; Popovska, N.; Holzmann, F.; Gerhard, H.; Topf, C.; Zenneck, U., [(1,5-Cyclooctadiene)(toluene)ruthenium(0)]: A Novel Precursor for the MOCVD of Thin Ruthenium Films. Chemical Vapor Deposition 2005, 11, (2), 99-105. [65] Schneider, A.; Popovska, N.; Jipa, I.; Atakan, B.; Siddiqi, M. A.; Siddiqui, R.; Zenneck, U., Minimizing the carbon content of thin ruthenium films by MOCVD precursor complex design and process control. Chemical Vapor Deposition 2007, 13, (8), 389-395. [66] Choi, J.; Choi, Y.; Hong, J.; Tian, H.; Roh, J.-S.; Kim, Y.; Chung, T.-M.; Woo Oh, Y.; Kim, Y.; Kim, C. G.; No, K., Composition and Electrical Properties of Metallic Ru Thin Films Deposited Using Ru(C6H6)(C6H8) Precursor. Japanese Journal of Applied Physics 2002, 41, (11B), 6852-6856. [67] Ruthenium Tetroxide and Other Ruthenium Compounds? (http:/ / www. springerlink. com/ content/ n265k571444pw788/ ) [68] INHALATION OF RADIONUCLIDES AND CARCINOGENESIS (http:/ / www. osti. gov/ energycitations/ product. biblio. jsp?osti_id=4142243)
• Los Alamos National Laboratory – Ruthenium (http:/ / periodic. lanl. gov/ elements/ 44. html)
External links • Nano-layer of ruthenium stabilizes magnetic sensors (http:/ / www. brightsurf. com/ news/ headlines/ 32014/ Nano-layer_of_ruthenium_stabilizes_magnetic_sensors. html) • WebElements.com – Ruthenium (http:/ / www. webelements. com/ webelements/ elements/ text/ Ru/ index. html) pnb:مینیھتور
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Article Sources and Contributors Ruthenium Source: http://en.wikipedia.org/w/index.php?oldid=308564898 Contributors: A2Kafir, Achim1999, Addshore, Adrian.benko, Ahoerstemeier, Alexandrov, AlimanRuna, Andres, Archimerged, Archiver, Arkuat, Avb, AxelBoldt, AyJay, Basement12, Bassbonerocks, Bearcat, Beetstra, Benbest, BlueEarth, Bogey97, Borislav Dopudja, Brian Huffman, Bryan Derksen, CYD, Cadmium, CanisRufus, Carabinieri, Carltzau, Carnildo, Chewyrunt, Chrislk02, Conversion script, Corpx, DDima, Dajwilkinson, Darrien, David Latapie, Delta G, DerHexer, Dilbert2000, Donarreiskoffer, DragonflySixtyseven, Drrocket, Dwmyers, Edgar181, Edhale, El C, Emperorbma, Epharos, Eudialytos, Femto, Francs2000, Frank Lofaro Jr., Globalistgirl, Goudzovski, GraemeL, GrahamHardy, Grendelkhan, Hadal, Hallpriest9, Hannibal, Harryboyles, Hede2000, Helge Skjeveland, Icairns, Ideyal, InnerJustice, Inter, IstvanWolf, Itub, Iulius, Ivan05, Jaraalbe, JasonAQuest, Joanjoc, John, Jose77, Kevlar67, Kia22, Klausok, Kubra, Kurykh, Kwamikagami, LA2, LinguisticDemographer, Mac, MathGuy, Maury Markowitz, Mav, Maximus Rex, Metahacker, Mgimpel, Minesweeper, Moondisaster, Mr0t1633, Mrshaba, Nergaal, Nick Y., Nihil novi, Ostap R, OverlordQ, Physchim62, Plantsurfer, Plasmic Physics, PlatinumX, Poindexter Propellerhead, Poolkris, Psyche825, Quintote, RSido, RTC, Remember, Reyk, Rifleman 82, Rkinch, Roberta F., Romanm, Rossnorman, Roybb95, Rursus, Saperaud, Schneelocke, Sengkang, Shaddack, Shafei, Sillybilly, Sl, Slashme, Smalljim, Socraticus, Socrplaya009, Spamhog, Spellage, Squids and Chips, Stifynsemons, Stone, Suisui, Sunborn, Svante, Tagishsimon, Tanada, Tassedethe, TerraFrost, Tetracube, Titus III, Tsorro, Untifler, V8rik, Vsmith, Watch37264, Yuniq, Yyy, Zahnrad, Zoicon5, 163 anonymous edits
Image Sources, Licenses and Contributors image:Ru-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Ru-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Georg Slickers, Mav, Paddy, Saperaud Image: Ruthenium powder.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Ruthenium_powder.jpg License: unknown Contributors: User:Materialscientist Image:Activity of pt group metals from uranium fission.png Source: http://en.wikipedia.org/w/index.php?title=File:Activity_of_pt_group_metals_from_uranium_fission.png License: Public Domain Contributors: Original uploader was Cadmium at en.wikipedia
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14
Rhodium
1
Rhodium 45
ruthenium ← rhodium → palladium
Co ↑
Rh ↓
Ir Periodic Table - Extended Periodic Table
General Name, Symbol, Number
rhodium, Rh, 45
Element category
transition metals
Group, Period, Block
9, 5, d
Appearance
silvery white metallic
Standard atomic weight
102.90550(2) g·mol
Electron configuration
[Kr] 4d 5s
−1
8
Electrons per shell
1
2, 8, 18, 16, 1 Physical properties
Phase
solid
Density (near r.t.)
12.41 g·cm
Liquid density at m.p.
10.7 g·cm−3
Melting point
2237 K (1964 °C, 3567 °F)
Boiling point
3968 K (3695 °C, 6683 °F)
Heat of fusion
26.59 kJ·mol−1
Heat of vaporization
494 kJ·mol−1
Specific heat capacity
(25 °C) 24.98 J·mol−1·K−1
−3
Vapor pressure P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
2288
2496
2749
3063
3405
3997
Atomic properties Crystal structure
cubic face centered
Rhodium
2 [1]
Oxidation states
6, 5, 4, 3, 2, 1 , -1 (amphoteric oxide)
Electronegativity
2.28 (Pauling scale)
Ionization energies
1st: 719.7 kJ/mol 2nd: 1740 kJ/mol 3rd: 2997 kJ/mol
Atomic radius
134 pm
Covalent radius
142±7 pm Miscellaneous
Magnetic ordering Electrical resistivity
[2]
paramagnetic
(0 °C) 43.3 nΩ·m
Thermal conductivity
(300 K) 150 W·m
Thermal expansion
(25 °C) 8.2 µm·m
Speed of sound (thin rod)
(20 °C) 4700 m/s
Young's modulus
380 GPa
Shear modulus
150 GPa
Bulk modulus
275 GPa
−1
−1
Poisson ratio
0.26
Mohs hardness
6.0
Vickers hardness
1246 MPa
Brinell hardness
1100 MPa
CAS registry number
·K
7440-16-6 Most-stable isotopes
−1
·K
−1
Rhodium
3
Main article: Isotopes of rhodium iso 99
Rh
NA syn
half-life 16.1 d
DM ε
-
γ
101m
Rh
101
Rh
syn
syn
4.34 d
3.3 y
Rh
syn
2.9 y
Rh
syn
207 d
Rh
105
Rh
100% syn
Ru
-
101
IT
0.157
101
γ
0.306, 0.545
ε
-
ε
ε
0.127, 0.198, 0.325 0.475, 0.631, 0.697, 1.046
Ru Rh
101
Ru
-
102
Ru
-
-
102
+
0.826, 1.301
102
β
−
1.151
102
γ
0.475, 0.628
β
103
99
-
γ
102
0.089, 0.353, 0.528
DP
ε
γ
102m
DE (MeV)
Ru Ru Pd
-
103
Rh is stable with 58 neutron
35.36 h
β
0.247, 0.260, 0.566
γ
0.306, 0.318
−
105
Pd
-
References
Rhodium (pronounced /ˈroʊdiəm/) is a chemical element that is a rare, silvery-white, hard transition metal and a member of the platinum group. Rhodium is found in platinum ores and is used in alloys with platinum and as a catalyst. It is abbreviated to Rh and has atomic number 45. It is one of the most expensive precious metals.
Characteristics Rhodium is a hard silvery white and durable metal that has a high reflectance. Rhodium metal does not normally form an oxide, even when heated.[3] Oxygen is absorbed from the atmosphere at the melting point of rhodium, but on solidification, the oxygen is released.[4] Rhodium has both a higher melting point and lower density than platinum. It is not attacked by acids: it is completely insoluble in nitric acid and dissolves slightly in aqua regia.
Rhodium
4
Chemical properties Rhodium belongs to group 9 in the periodic table, Z
Element
No. of electrons/shell
27
cobalt
2, 8, 15, 2
45
rhodium
2, 8, 18, 16, 1
77
iridium
2, 8, 18, 32, 15, 2
109
meitnerium
2, 8, 18, 32, 32, 15, 2
but has an atypical configuration in its outermost electron shells compared to the rest of the members. (This can also be observed in the neighborhood of niobium (41), ruthenium (44), rhodium (45), and palladium (46).) Oxidation states of rhodium +0
Rh4(CO)12
+1
RhCl(PH3)2
+2
Rh2(O2CCH3)
+3
RhCl3, Rh2O3
+4
RhF4, RhO2
+5
RhF5, Sr3LiRhO6
+6
RhF6
Common oxidation states of rhodium is +3, but oxidation states from +0 to +6 are observed.[5] Unlike ruthenium and osmium, rhodium forms no volatile oxygen compounds. The known stable oxides include Rh2O3, RhO2, RhO2·xH2O, Na2RhO3, Sr3LiRhO6 and Sr3NaRhO6[6] . Halogen compounds are known in nearly the full range of possible oxidation states. Rhodium(III) chloride, rhodium(IV) fluoride, rhodium(V) fluoride and rhodium(VI) fluoride are some examples. The lower oxidation states are only stable if ligands are present.[7] The best known example is the Wilkinson's catalyst chlorotris(triphenylphosphine)rhodium(I). The catalyst is for example used for the hydrogenation of alkenes.[8]
Isotopes Naturally occurring rhodium is composed of only one isotope, 103Rh. The most stable radioisotopes are 101Rh with a half-life of 3.3 years, 102Rh with a half-life of 207 days, 102mRh with a half-life of 2.9 years, and 99Rh with a half-life of 16.1 days. Twenty other radioisotopes have been characterized with atomic weights ranging from 92.926 u (93Rh) to 116.925 u (117Rh). Most of these have half-lives that are less than an hour except 100Rh (half-life: 20.8 hours) and 105 Rh (half-life: 35.36 hours). There are also numerous meta states with the most stable being 102mRh (0.141 MeV) with a half-life of about 2.9 years and 101mRh (0.157 MeV) with a half-life of 4.34 days. See isotopes of rhodium.[9] Wilkinson's catalyst
Rhodium
5
The primary decay mode before the only stable isotope, 103Rh, is electron capture and the primary mode after is beta emission. The primary decay product before 103Rh is ruthenium and the primary product after is palladium.[10]
History Rhodium (Greek rhodon meaning "rose") was discovered in 1803 by William Hyde Wollaston,[11] [12] soon after his discovery of palladium.[13] [14] He made this discovery in England using crude platinum ore that he presumably obtained from South America.[15] His procedure involved dissolving the ore in aqua regia and neutralizing the acid with sodium hydroxide (NaOH). He then precipitated the platinum by adding ammonium chloride, NH4Cl, as ammonium chloroplatinate. All other metals like copper, lead, palladium and rhodium were precipitated with zinc. Diluted nitric acid dissolved all but palladium and rhodium, which were dissolved in aqua regia and the rhodium was precipitated by the addition of sodium chloride as Na3[RhCl6]·nH2O. After washing with ethanol, the rose red precipitate was reacted with zinc forming rhodium metal.[16]
Applications The primary use of this element is in automobiles as a catalytic converter, which converts harmful emissions from the engine into less harmful gases.[17] [18]
Rhodium foil and wire
Catalytic converter
Cross section of a Metal-core Converter
In 2007 81%[17] of the world production of rhodium was consumed to produce three-way catalytic converters.[17] Rhodium shows some advantages over the other platinum metals in the reduction of nitrogen oxides to nitrogen and oxygen:[19] 2NOx → xO2 + N2
The recycling of catalytic converters also became a valuable source for rhodium. In 2007 5.7 t were extracted from this source. Compared to the 22 t which have been mined, this is a relatively high recycling rate.[17]
Rhodium
6
Other uses Rhodium is used as an alloying agent for hardening and improving the corrosion resistance[3] of platinum and palladium. These alloys are used in furnace windings, bushings for glass fiber production, thermocouple elements, electrodes for aircraft spark plugs, and laboratory crucibles.[20] Other uses include: • An electrical contact material due to its low electrical resistance, low and stable contact resistance, and high corrosion resistance.[21] • Plated rhodium, made by electroplating or evaporation, is extremely hard and is used for optical instruments.[22]
Rhodium plated white gold wedding ring
• This metal finds use in jewelry and for decorations. It is electroplated on white gold and platinum to give it a reflective white surface. This is known as rhodium flashing in the jewelry business. It also may be used in coating sterling silver in order to strengthen the metal from tarnish, as a result from the copper compound found in sterling silver. Solid (pure) Rhodium jewelry is very rare, because the metal has both high melting point and poor malleability (making such jewelry very hard to fabricate) rather than due to its high price.[23]
• It is also a highly useful catalyst in a number of industrial processes. Notably, it is used in the automobile catalytic converters and for catalytic carbonylation of methanol to produce acetic acid by the Monsanto process)[24] It is also used to catalyze addition of hydrosilanes to molecular double bonds, a process important in manufacture of certain silicone rubbers.[25] • Rhodium catalysts can be used to reduce benzene to cyclohexane.[26] • The complex of a rhodium ion with BINAP gives a widely used chiral catalyst for chiral synthesis, as in the synthesis of menthol.[27] • It is also used as a filter in mammography systems because of the characteristic X-rays it produces.[28] • It is also used in high quality pen surfaces due to its high chemical and mechanical resistance. These pens include Graf von Faber-Castell[29] and Caran D'ache[30] . • Rhodium neutron detectors are used in Combustion Engineering Nuclear Reactors to measure neutron flux levels - a method that requires a digital filter to determine the current neutron flux level, as there are three signals generated: immediate, a few seconds later, and a minute later, each with its own signal level, and all three are combined in the rhodium detector signals. The three Palo Verde nuclear reactors each have 305 rhodium neutron detectors, 61 detectors on each of 5 vertical levels, providing an accurate 3-D "picture" of reactivity, allowing fine tuning to most economically burn the nuclear fuel.[31]
Occurrence Normal mining The industrial extraction of rhodium is complex as the metal occurs in ores mixed with other metals such as palladium, silver, platinum, and gold. It is found in platinum ores and obtained free as a white inert metal which is very difficult to fuse. Principal sources of this element are located in South Africa, in river sands of the Ural Mountains, and in North
Rhodium
7
America, including the copper-nickel sulfide mining area of the Sudbury, Ontario region. Although the quantity at Sudbury is very small, the large amount of processed nickel ore makes rhodium recovery cost effective. The main exporter of rhodium is South Africa (>80%) followed by Russia.[32] The annual world production of this element is only about 25 tons and there are very few rhodium-bearing minerals. As of October 2007, rhodium costs approximately eight times more than gold, 450 times more than silver, and 27,250 times more than copper by weight. Rhodium's typical historical price is about $1,000/troy oz,[33] but in recent years, it has increased to about $4500/troy oz.[34] In 2008 the price briefly rose above $10,000 per ounce.[34] The 3rd quarter 2008 economic slowdown has pushed prices sharply back below $1,000 per ounce, however.[34]
Fission product It is also possible to extract rhodium from used nuclear fuel, which contains rhodium (1 kg of the fission products of 235U contains 13.3 grams of 103Rh). As a typical used fuel has 3% fission products by weight, it will contain about 400 grams of rhodium per ton of used fuel. The longest lived radioisotope of rhodium is 102mRh which has a half life of 2.9 years, whereas the ground state (102Rh) has a half life of 207 days. One
kilogram
of
fission
rhodium will contain 6.62 ng of 102Rh and 3.68 ng of 102m Rh. As 102Rh decays by beta decay to either 102Ru (80%) (some positron emission will occur) or 102Pd (20%) (gamma ray photons with about 500 keV are generated) and the excited state decays by beta decay (electron 102 capture) to Ru (gamma ray photons with about 1 MeV are The radioactivity in MBq per gram of each of the platinum group generated). If the fission metals which are formed by the fission of uranium, ruthenium is the most radioactive. Palladium has an almost constant activity due to the occurs in an instant then 13.3 very long lived 107Pd, while rhodium is the least radioactive grams of rhodium will contain 67.1 MBq (1.81 mCi) of 102Rh and 10.8 MBq (291 μCi) of 102mRh. As it is normal to allow used nuclear fuel to rest for about five years before reprocessing, much of this activity will decay leaving 4.7 MBq of 102 Rh and 5.0 MBq of 102mRh. If the rhodium metal was then left for 20 years after fission, then the 13.3 grams of rhodium metal would contain 1.3 kBq of 102Rh and 500 kBq of 102m Rh. At first glance, the rhodium might be adding to the resource value of reprocessed fission waste, but the cost of the separation of rhodium from other metals needs to be considered.[35]
Rhodium
8
Precautions Rhodium metal is, as a noble metal, inert. However, when rhodium is chemically bound, it is reactive. Lethal intake (LD50) for rats is 12.6 mg/kg of rhodium chloride (RhCl3).[36] Rhodium compounds can strongly stain human skin. The element plays no biological role in humans. If used in elemental form rather than as compounds, the metal is harmless.[37]
Ornamental uses Rhodium has been used for honours, or to symbolize wealth, when more commonly used metals such as silver, gold, or platinum are deemed insufficient. In 1979 the Guinness Book of World Records gave Paul McCartney a rhodium-plated disc for being history's all-time best-selling songwriter and recording artist.[38]
See also • Rhodium compounds
External links • • • •
WebElements.com – Rhodium [39] Current Rhodium price [40] Rhodium Technical and Safety Data [41] Los Alamos National Laboratory – Rhodium
[42]
References [1] " Rhodium: rhodium(I) fluoride compound data (http:/ / openmopac. net/ data_normal/ rhfr_jmol. html)". OpenMOPAC.net. . Retrieved 2007-12-10. [2] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [3] Cramer, Stephen; S., Jr Covino, Bernard (1990). ASM handbook (http:/ / books. google. com/ books?id=QV0sWU2qF5oC& pg=PA396). Materials Park, OH: ASM International. pp. 393–396. ISBN 0-87170-707-1. . [4] Emsley, John (2001). Nature's Building Blocks ((Hardcover, First Edition) ed.). Oxford University Press. pp. 363. ISBN 0198503407. [5] Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils; (1985). Lehrbuch der Anorganischen Chemie (91–100 ed.). Walter de Gruyter. pp. 1056–1057. ISBN 3-11-007511-3. [6] Reisner, B. A.; Stacy, A. M. (1998). 120. pp. 9682–9989. doi: 10.1021/ja974231q (http:/ / dx. doi. org/ 10. 1021/ ja974231q). [7] Griffith, W. P. The Rarer Platinum Metals; John Wiley and Sons: New York, 1976; p 313. [8] Osborn, J. A.; Jardine, F. H.; Young, J. F.; Wilkinson, G. (1966). "The Preparation and Properties of Tris(triphenylphosphine)halogenorhodium(I) and Some Reactions Thereof Including Catalytic Homogeneous Hydrogenation of Olefins and Acetylenes and Their Derivatives". Journal of the Chemical Society A: 1711–1732. doi: 10.1039/J19660001711 (http:/ / dx. doi. org/ 10. 1039/ J19660001711). [9] Audi, G. (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3–128. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [10] David R. Lide (ed.), Norman E. Holden in CRC Handbook of Chemistry and Physics, 85th Edition CRC Press. Boca Raton, Florida (2005). Section 11, Table of the Isotopes. [11] " WebElements - The History of Rhodium (http:/ / www. webelements. com/ webelements/ elements/ text/ Rh/ hist. html)". . Retrieved 2009-02-06. [12] Wollaston, W. H. (1805). "On the Discovery of Palladium; With Observations on Other Substances Found with Platina". Philosophical Transactions of the Royal Society of London 95: 316–330. doi: 10.1098/rstl.1805.0024 (http:/ / dx. doi. org/ 10. 1098/ rstl. 1805. 0024).
Rhodium [13] W. P. Griffith (2003). " Rhodium and Palladium - Events Surrounding Its Discovery (http:/ / www. platinummetalsreview. com/ dynamic/ article/ view/ 47-4-175-183)". Platinum Metals Review 47 (4): 175–183. . [14] Wollaston, W. H. (1804). "On a New Metal, Found in Crude Platina". Philosophical Transactions of the Royal Society of London 94: 419–430. doi: 10.1098/rstl.1804.0019 (http:/ / dx. doi. org/ 10. 1098/ rstl. 1804. 0019). [15] Lide, David R (2004). CRC handbook of chemistry and physics: a ready-reference book of chemical and physical data. Boca Raton: CRC Press. pp. 4–26. ISBN 0-8493-0485-7. [16] Griffith, W. P. (2003). "Bicentenary of Four Platinum Group Metals: Osmium and iridium – events surrounding their discoveries". Platinum Metals Review 47 (4): 175–183. [17] George, Micheal W.. " Commodity Report: Platinum-Group Metals (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ platinum/ mcs-2008-plati. pdf)". United States Geological Survey USGS. . Retrieved 2008-09-16. [18] George, Micheal W.. " 2006 Minerals Yearbook: Platinum-Group Metals (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ platinum/ myb1-2006-plati. pdf)". United States Geological Survey USGS. . Retrieved 2008-09-16. [19] Shelef, M.; Graham, G. W. (1994). "Why Rhodium in Automotive Three-Way Catalysts?". Catalysis Reviews 36 (3): 433–457. doi: 10.1080/01614949408009468 (http:/ / dx. doi. org/ 10. 1080/ 01614949408009468). [20] Lide, David R (2004). CRC handbook of chemistry and physics: a ready-reference book of chemical and physical data (http:/ / books. google. com/ books?id=WDll8hA006AC). Boca Raton: CRC Press. pp. 4–26. ISBN 0-8493-0485-7. . [21] Weisberg, Alfred M. (1999). "Rhodium plating". Metal Finishing 97 (1): 296–299. doi: 10.1016/S0026-0576(00)83088-3 (http:/ / dx. doi. org/ 10. 1016/ S0026-0576(00)83088-3). [22] Smith, Warren J. (2007). " Reflectors (http:/ / books. google. de/ books?id=DrtM_bAnf_YC)". Modern optical engineering: the design of optical systems. McGraw-Hill. pp. 247–248. ISBN 9780071476874. . [23] Fischer, Torkel (1984). "Contact sensitivity to nickel in white gold". Contact Dermatitis 10: 23–24. doi: 10.1111/j.1600-0536.1984.tb00056.x (http:/ / dx. doi. org/ 10. 1111/ j. 1600-0536. 1984. tb00056. x). [24] Roth, James F. (1975). " Rhodium Catalysed Carbonylation of Methanol (http:/ / www. platinummetalsreview. com/ pdf/ pmr-v19-i1-012-014. pdf)" (PDF). Platinum Metals Review 19 (1 January): 12–14. . [25] M. Heidingsfeldova and M. Capka "Rhodium complexes as catalysts for hydrosilylation crosslinking of silicone rubber" Journal of Applied Polymer Science 30 (2003) 1837 (http:/ / dx. doi. org/ 10. 1002/ app. 1985. 070300505) [26] S. B. Halligudi et al. "Hydrogenation of benzene to cyclohexane catalyzed by rhodium(I) complex supported on montmorillonite clay" Reaction Kinetics and Catalysis Letters 48 (1992) 547 (http:/ / dx. doi. org/ 10. 1007/ BF02162706) [27] S. Akutagawa "Asymmetric synthesis by metal BINAP catalysts" Applied Catalysis A: 128 (1995) 171 (http:/ / dx. doi. org/ 10. 1016/ 0926-860X(95)00097-6) [28] C P McDonagh, J L Leake and S A Beaman "Optimum x-ray spectra for mammography: choice of K-edge filters for tungsten anode tubes" Phys. Med. Biol. 29 (1984) 249 (http:/ / dx. doi. org/ 10. 1088/ 0031-9155/ 29/ 3/ 004) [29] Guilloche luxury pen range by Graf von Faber-Castell (http:/ / www. pensfromheaven. com/ pens_luxury_Graf-von-Faber-Castell_Guilloche. htm) [30] Caran D'Ache Ecridor Type 55 Rhodium Fountain Pen (http:/ / www. pengallery. com/ details. aspx?productID=2990) [31] cite journal | first = A. P. | last = Sokolov Pochivalin, G. P.; Shipovskikh, Yu. M.; Garusov, Yu. V.; Chernikov O. G.; Shevchenko V. G. | title = Rhodium self-powered detector for monitoring neutron fluence, energy production, and isotopic composition of fuel | doi = 10.1007/BF00844622 | journal = Atomic Energy | volume = 74 | year = 1993 pages = 365–367 [32] Chevalier, Patrick. " Mineral Yearbook: Platinum Group Metals (http:/ / nrcan. gc. ca/ mms/ cmy/ content/ 2004/ 71. pdf)". Natural Resources Canada. . Retrieved 2008-10-17. [33] (http:/ / periodic. lanl. gov/ elements/ 45. html) [34] KITCO Rhodium Price Charts (http:/ / www. kitco. com/ charts/ rhodium. html) [35] Bush, R. P. (1991). "Recovery of Platinum Group Metals from High Level Radioactive Waste". Platinum Metals Review 35 (4): 202–208. [36] Landolt, Robert R.; Berk Harold W.; Russell, Henry T. (1972). "Studies on the toxicity of rhodium trichloride in rats and rabbits". Toxicology and Applied Pharmacology 21 (4): 589–590. doi: 10.1016/0041-008X(72)90016-6 (http:/ / dx. doi. org/ 10. 1016/ 0041-008X(72)90016-6). [37] Leikin, Jerrold B.; Paloucek Frank P. (2008). Poisoning and Toxicology Handbook (http:/ / books. google. com/ books?id=0Bw2UJTC_uMC). Informa Health Care. pp. 846. ISBN 9781420044799. . [38] " Hit & Run: Ring the changes (http:/ / www. independent. co. uk/ news/ people/ hit-and-run/ hit--run-ring-the-changes-1044166. html)". The Independent. . Retrieved 2009-06-06.
9
Rhodium [39] [40] [41] [42]
http:/ / www. webelements. com/ webelements/ elements/ text/ Rh/ index. html http:/ / www. kitco. com/ market/ http:/ / www. americanelements. com/ rh. html http:/ / periodic. lanl. gov/ elements/ 45. html
10
Article Sources and Contributors
Article Sources and Contributors Rhodium Source: http://en.wikipedia.org/w/index.php?oldid=307293521 Contributors: Achaemenes, Achim1999, Agathoclea, Ahoerstemeier, Aitias, Akamad, Alansohn, AlimanRuna, Andrewpmk, Archimerged, Arkuat, B07, Babakathy, Bearcat, Beetstra, Benbest, Betavenator, Bipedal, Blanchardb, BlueEarth, Borislav Dopudja, Brandonrush, BrokenSegue, Bryan Derksen, Bryancpark, CYD, Cadmium, Carnildo, Chrissant, Closedmouth, Conversion script, Cool Blue, Dantheman531, Darrien, David Latapie, Delta G, Deor, Dirty apes, Donarreiskoffer, DrKranium, Dschwen, Dson72791, Eaolson, Edgar181, El C, Ellmist, Emok, Emperorbma, Epbr123, Ertdredge, Femto, Frank Lofaro Jr., Friendly Neighbour, Gaius Cornelius, Garsanllean, Glengalbraith, Gmcole, Gogo Dodo, Grandor, Grendelkhan, Gringer, Gurch, H.h.kan, Hak-kâ-ngìn, Hallpriest9, Hede2000, Hermitian, Hipocrite, Hyperdeath, Icairns, Ideyal, Idotperson, Iridescent, Ixfd64, J.delanoy, JForget, JWB, Jaraalbe, Jeffrey Vernon Merkey, Jennyvu96, Jimp, Joanjoc, John, Jonathan Watt, Jose77, Kachyna, Katalaveno, Katieh5584, Kent Wang, Ktsquare, Kumorifox, Kwamikagami, Lady BlahDeBlah, Ladyceltic343, Lankiveil, LarryMorseDCOhio, Lauren 218, LionFaceMan, LuisVilla, Madmardigan53, Markjoseph125, Materialscientist, Mav, Mdf, Mervyn, Mgimpel, Michael Schubart, Minesweeper, Morandavid575, Mortdefides, N2e, Naddy, Nergaal, Nihiltres, Nishkid64, Paddu, Patstuart, PaulHanson, Physchim62, Plasmic Physics, PlatinumX, Polimerek, Poolkris, Pras, Publunch, Qwasty, Qwfp, RSido, RTC, Rbbwiki, Remember, Rhodium, Richardmilgate, Ritterrat, Roadahead, Romanm, Rossnorman, ST47, Saga City, Saperaud, Savant13, Sbharris, Schneelocke, Schrodingers rabbit, Sengkang, Shaddack, Shafei, Shanqz, Skatebiker, Skidude9950, Sl, Sleigh, Squids and Chips, Stifynsemons, Stone, Suisui, Sunborn, Tagishsimon, Tetracube, The Rambling Man, Titus III, TomasBat, Tsorro, Vatic7, Vectro, Vercingetorix08, Vivaldi, Voortle, Voyagerfan5761, Vsmith, Vuo, Waggers, Walkerma, Warut, Wasted Time R, Watch37264, Wikidemon, XIDE, Yekrats, Zfr, 267 anonymous edits
Image Sources, Licenses and Contributors image:Rh-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Rh-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Georg Slickers, Mav, Paddy, Saperaud Image: Rh,45.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Rh,45.jpg License: GNU Free Documentation License Contributors: User:RTC Image:Wilkinson's-catalyst-2D.png Source: http://en.wikipedia.org/w/index.php?title=File:Wilkinson's-catalyst-2D.png License: Public Domain Contributors: Benjah-bmm27 Image:Rhodium foil and wire.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Rhodium_foil_and_wire.jpg License: unknown Contributors: User:Dschwen Image:Aufgeschnittener Metall Katalysator für ein Auto.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Aufgeschnittener_Metall_Katalysator_für_ein_Auto.jpg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User Stahlkocher on de.wikipedia Image:White-gold--rhodium-plated.jpg Source: http://en.wikipedia.org/w/index.php?title=File:White-gold--rhodium-plated.jpg License: unknown Contributors: User:Schtone Image:Activity of pt group metals from uranium fission.png Source: http://en.wikipedia.org/w/index.php?title=File:Activity_of_pt_group_metals_from_uranium_fission.png License: Public Domain Contributors: Original uploader was Cadmium at en.wikipedia
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11
Palladium
1
Palladium 46
rhodium ← palladium → silver
Ni ↑
Pd ↓
Pt Periodic Table - Extended Periodic Table
General Name, Symbol, Number
palladium, Pd, 46
Element category
transition metals
Group, Period, Block
10, 5, d
Appearance
silvery white metallic
Standard atomic weight
106.42(1) g·mol
Electron configuration
[Kr] 4d
−1
10
Electrons per shell
2, 8, 18, 18, 0 Physical properties
Phase
solid
Density (near r.t.)
12.023 g·cm−3
Liquid density at m.p.
10.38 g·cm−3
Melting point
1828.05 K (1554.9 °C, 2830.82 °F)
Boiling point
3236 K (2963 °C, 5365 °F)
Heat of fusion
16.74 kJ·mol−1
Heat of vaporization
362 kJ·mol−1
Specific heat capacity
(25 °C) 25.98 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1721
1897
2117
2395
2753
3234
Atomic properties
Palladium
2
Crystal structure
cubic face centered
Oxidation states
0, +1, +2, +4, +6 (mildly basic oxide)
Electronegativity Ionization energies
2.20 (Pauling scale) 1st: 804.4 kJ/mol 2nd: 1870 kJ/mol 3rd: 3177 kJ/mol
Atomic radius
137 pm
Covalent radius
139±6 pm
Van der Waals radius
163 pm Miscellaneous [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(20 °C) 105.4 n Ω·m
Thermal conductivity
(300 K) 71.8 W·m
Thermal expansion
(25 °C) 11.8 µm·m
Speed of sound (thin rod)
(20 °C) 3070 m/s
Young's modulus
121 GPa
Shear modulus
44 GPa
Bulk modulus
180 GPa
−1
−1
Poisson ratio
0.39
Mohs hardness
4.75
Vickers hardness
461 MPa
Brinell hardness
37.3 MPa
CAS registry number
7440-05-3 Most-stable isotopes
−1
·K
−1
·K
Palladium
3
Main article: Isotopes of palladium iso 100
Pd
NA syn
half-life 3.63 d
DM ε
-
γ
102
Pd
103
Pd
1.02% syn
16.991 d 104
105
22.33%
105
106
27.33%
106
Pd
107
Pd
syn
-
ε
-
103
Rh
Pd is stable with 59 neutron Pd is stable with 60 neutron 6
6.5×10 y
26.46%
108
110
11.72%
110
Pd
Rh
Pd is stable with 58 neutron
108
Pd
100
Pd is stable with 56 neutron
11.14%
Pd
0.084, 0.074, 0.126
DP
102
104
Pd
DE (MeV)
−
β
0.033
107
Ag
Pd is stable with 62 neutron Pd is stable with 64 neutron References
Palladium (pronounced /pəˈleɪdiəm/) is a chemical element with the chemical symbol Pd and an atomic number of 46. Palladium is a rare and lustrous silvery-white metal that was discovered in 1803 by William Hyde Wollaston, who named it after the asteroid Pallas, which in turn, was named after the epithet of the Greek goddess Athena, acquired by her when she slew Pallas. Palladium, along with platinum, rhodium, ruthenium, iridium and osmium form a group of elements referred to as the platinum group metals (PGMs). Platinum group metals share similar chemical properties, but palladium has the lowest melting point and is the least dense of these precious metals. The unique properties of palladium and other platinum group metals account for their widespread use. One in four goods manufactured today either contain platinum group metals or had platinum group metals play a key role during their manufacturing process[2] . Over half of the supply of palladium and its congener platinum goes into catalytic converters, which convert up to 90% of harmful gases from auto exhaust (hydrocarbons, carbon monoxide and nitrogen oxide) into less harmful substances (nitrogen, carbon dioxide and water vapor). Palladium is found in many electronics including computers, mobile phones, multi-layer ceramic capacitors, component plating, low voltage electrical contacts, and SED/OLED/LCD televisions. Palladium is also used in dentistry, medicine, hydrogen purification, chemical applications, and groundwater treatment. Palladium plays a key role in the technology used for fuel cells, which combines hydrogen and oxygen to produce electricity, heat and water. Palladium bullion has ISO currency codes of XPD and 964. Palladium is one of only four metals to have such codes, the others being gold, silver and platinum. Ore deposits of palladium and other platinum group metals are rare, and the most extensive deposits have been found in the norite belt of the Bushveld Igneous Complex in the
Palladium
4
Transvaal in South Africa, the Stillwater Complex in Montana, United States, the Sudbury District of Ontario, Canada, and the Norilsk Complex in Russia. In addition to mining, recycling is also a source of palladium, mostly from scrapped catalytic converters. The numerous applications and limited supply sources of palladium result in palladium drawing considerable investment interest.
History Palladium was discovered by William Hyde Wollaston in 1803.[3] [4] This element was named by Wollaston in 1804 after the asteroid Pallas, which had been discovered two years earlier.[5] Wollaston found palladium in crude platinum ore from South America by dissolving the ore in aqua regia, neutralizing the solution with sodium hydroxide, and precipitating platinum as ammonium chloroplatinate with ammonium chloride. He added mercuric cyanide to form the compound palladium cyanide, which was heated to extract palladium metal. Palladium chloride was at one time prescribed as a tuberculosis treatment at the rate of 0.065g per day (approximately one milligram per kilogram of body weight). This treatment did have many negative side-effects, and was later replaced by more effective drugs.[6] Palladium's affinity for hydrogen led it to play an essential role in the Fleischmann-Pons experiment in 1989, also known as cold fusion. In the run up to 2000, Russian supply of palladium to the global market was repeatedly delayed and disrupted[7] because the export quota was not granted on time, for political reasons. The ensuing market panic drove the palladium price to an all-time high of $1100 per ounce in January 2001.[8] Around this time, the Ford Motor Company, fearing auto vehicle production disruption due to a possible palladium shortage, stockpiled large amounts of the metal purchased near the price high. When prices fell in early 2001, Ford lost nearly US$1 billion.[9] World demand for palladium increased from 100 tons in 1990 to nearly 300 tons in 2000. The global production of palladium from mines was 222 metric tons in 2006 according to USGS data.[10] Most palladium is used for catalytic converters in the automobile industry.[11]
Occurrence In 2005, Russia was the top producer of palladium, with at least 50% world share, followed by South Africa, Canada and the U.S., reports the British Geological Survey. Palladium may be found as a free metal alloyed with gold and other platinum group metals in placer deposits Palladium output in 2005 of the Ural Mountains, Australia, Ethiopia, South and North America. It is commercially produced from nickel-copper deposits found in South Africa, Ontario, and Siberia; It takes processing of many metric tons of ore to extract just one troy ounce of palladium. However, the mine production could still be profitable, depending on current metal prices, as other metals are produced together: nickel, copper, platinum and rhodium. The world's largest single producer of palladium is MMC Norilsk Nickel produced from the Norilsk–Talnakh nickel deposits. The Merensky Reef of the Bushveld Igneous Complex of
Palladium
5
South Africa contains significant palladium in addition to other platinum group elements. The Stillwater igneous complex of Montana and the Roby zone orebody of the Lac des Îles igneous complex of Ontario also contain mineable palladium. Palladium is also produced in nuclear fission reactors and can be extracted from spent nuclear fuel (see synthesis of noble metals) though the quantity produced is insignificant. Palladium is found in the rare minerals cooperite and polarite.
Characteristics Palladium belongs to group 10 in the periodic table: Z
Element
No. of electrons/shell
28
nickel
2, 8, 16, 2
46
palladium
2, 8, 18, 18
78
platinum
2, 8, 18, 32, 17, 1
110
darmstadtium
2, 8, 18, 32, 32, 17, 1
but has a very atypical configuration in its outermost electron shells compared to the rest of the members of group 10, if not to all elements. (See also niobium (41), ruthenium (44), and rhodium (45).) Palladium is a soft silver-white metal that resembles platinum. It is the least dense and has the lowest melting point of the platinum group metals. It is soft and ductile when annealed and greatly increases its strength and hardness when it is cold-worked. Palladium dissolves slowly in sulfuric, nitric, and hydrochloric acid.[5] This metal also does not react with oxygen at normal temperatures (and thus does not tarnish in air). Palladium heated to 800°C will produce a layer of palladium(II) oxide (PdO). It lightly tarnishes in moist atmosphere containing sulfur. The metal has the uncommon ability to absorb up to 900 times its own volume of hydrogen at room temperatures. It is thought that this possibly forms palladium hydride (PdH2) but it is not yet clear if this is a true chemical compound.[5] When palladium has absorbed large amounts of hydrogen, it will expand slightly in size.[12] Common oxidation states of palladium are 0,+1, +2 and +4. Although originally +3 was thought of as one of the fundamental oxidation states of palladium, there is no evidence for palladium occurring in the +3 oxidation state; this has been investigated via X-ray diffraction for a number of compounds, indicating a dimer of palladium(II) and palladium(IV) instead. Recently, compounds with an oxidation state of +6 were synthesised.
Isotopes Naturally-occurring palladium is composed of six isotopes. The most stable radioisotopes are 107Pd with a half-life of 6.5 million years, 103Pd with a half-life of 17 days, and 100Pd with a half-life of 3.63 days. Eighteen other radioisotopes have been characterized with atomic weights ranging from 92.936 u (93Pd) to 119.924 u (120Pd). Most of these have half-lives that are less than a half-hour, except 101Pd (half-life: 8.47 hours), 109Pd (half-life: 13.7 hours), and 112Pd (half-life: 21 hours). The primary decay mode before the most abundant stable isotope, 106Pd, is electron capture and the primary mode after is beta decay. The primary decay product before 106Pd
Palladium is rhodium and the primary product after is silver. Radiogenic 107Ag is a decay product of 107Pd and was first discovered in the Santa Clara, California meteorite of 1978.[13] The discoverers suggest that the coalescence and differentiation of iron-cored small planets may have occurred 10 million years after a nucleosynthetic event. 107Pd versus Ag correlations observed in bodies, which have clearly been melted since accretion of the solar system, must reflect the presence of short-lived nuclides in the early solar system.[14]
Compounds Palladium chloride, bromide and acetate are reactive and relatively inexpensive, making them convenient entry points to palladium chemistry. All three compounds are not monomeric; the chloride and bromide often need to be refluxed in acetonitrile to obtain the more reactive acetonitrile complex monomers, e.g.:[15] PdCl2 + 2 MeCN → PdCl2(MeCN)2 The great many reactions in which palladium compounds serve as catalysts are collectively known as palladium coupling reactions. Prominent examples include the Heck, Suzuki reaction, and Stille reactions. Palladium(II) acetate, tetrakis(triphenylphosphine)palladium(0) (Pd(PPh3), and tris(dibenzylideneacetone)dipalladium(0) (Pd2(dba)3) are useful in this regard, either as catalysts, or as starting points to catalysts.
Applications Palladium is used in jewelry, in dentistry,[16] [17] watch making, in blood sugar test strips, in aircraft spark plugs and in the production of surgical instruments and electrical contacts.[18] Palladium is also used to make professional transverse flutes.[19]
Electronics The second biggest application of palladium in electronics is making the multilayer ceramic capacitor.[20] Palladium (and palladium-silver alloys) are used as electrodes in multi-layer ceramic capacitors.[16] Palladium (sometimes alloyed with nickel) is used in connector platings in consumer electronics. It is also used in plating of electronic components and in soldering materials. The electronic sector consumed 1.07 million troy ounces (33.2 tonnes) of palladium in 2006, according to a Johnson Matthey report.[21]
Technology Hydrogen easily diffuses through heated palladium; thus, it provides a means of purifying the gas.[5] Membrane reactors with Pd membranes are therefore used for the production of hydrogen. It is a part of the palladium-hydrogen electrode in electrochemical studies. Palladium(II) chloride can absorb large amounts of carbon monoxide gas, and is used in carbon monoxide detectors.
6
Palladium
7
Catalysis When it is finely divided, such as in palladium on carbon, palladium forms a good catalyst and is used to speed up hydrogenation and dehydrogenation reactions, as well as in petroleum cracking. A large number of carbon-carbon bond forming reactions in organic chemistry (such as the Heck and Suzuki coupling) are facilitated by catalysis with palladium compounds. The largest use of palladium today is in catalytic converters.[16] . In addition palladium, when dispersed on conductive materials, proves to be an excellent electrocatalyst for oxidation of primary alcohols in alkaline media.[22] Pd is also a versatile metal for homogeneous catalysis. It is used in combination with a broad variety of ligands for highly selective chemical transformations. A 2008 study showed that palladium is an effective catalyst for making carbon-fluoride bonds.[23]
Hydrogen storage Palladium hydride is metallic palladium that contains a substantial quantity of hydrogen within its crystal lattice. At room temperature and atmospheric pressure, palladium can absorb up to 900 times its own volume of hydrogen in a reversible process. This property has been investigated because hydrogen storage is of such interest and a better understanding of what happens at the molecular level could give clues to designing improved metal hydrides. A palladium based store, however, would be prohibitively expensive due to the cost of the metal.[24]
Jewelry Palladium itself has been used as a precious metal in jewelry since 1939, as an alternative to platinum or white gold. This is due to its naturally white properties, giving it no need for rhodium plating. It is slightly whiter, much lighter and about 12% harder than platinum. Similar to gold, palladium can be beaten into a thin leaf form as thin as 100 nm (1/250,000 in).[5] Like platinum, it will develop a hazy patina over time. Unlike platinum, however, palladium may discolor at high soldering temperatures, become brittle with repeated heating and cooling, and react with strong acids.[25]
A Palladium plated belt buckle.
Palladium is one of the three most popular metals used to make white gold alloys.[16] (Nickel and silver can also be used.) Palladium-gold is a more expensive alloy than nickel-gold, but seldom causes allergic reactions (though certain cross-allergies with nickel may occur).[26] When platinum was declared a strategic government resource during World War II, many jewelry bands were made out of palladium.[27] As recently as September 2001,[28] palladium was more expensive than platinum and rarely used in jewelry also due to the technical obstacle of casting. However the casting problem has been resolved, and its use in jewelry has increased because of a large spike in the price of platinum and a drop in the price of palladium.[29] Prior to 2004, the principal use of palladium in jewelry was as an alloy in the manufacture of white gold jewelry, but, beginning early in 2004 when gold and platinum prices began to
Palladium rise steeply, Chinese jewelers began fabricating significant volumes of palladium jewelry. Johnson Matthey estimated that in 2004, with the introduction of palladium jewelry in China, demand for palladium for jewelry fabrication was 920,000 ounces, or approximately 14% of the total palladium demand for 2004 - an increase of almost 700,000 ounces from the previous year. This growth continued during 2005, with estimated worldwide jewelry demand for palladium of about 1.4 million ounces, or almost 21% of net palladium supply, again with most of the demand centered in China. The popularity of palladium jewelry is expected to grow in 2008 as the world's biggest producers embark on a joint marketing effort to promote palladium jewelry worldwide [30]
Photography With the platinotype printing process photographers make fine-art black-and-white prints using platinum or palladium salts. Often used with platinum, palladium provides an alternative to silver.[31]
Art Palladium leaf is one of several alternatives to silver leaf used in manuscript illumination. The use of silver leaf is problematic due to its surprisingly fast tarnishing process. Aluminum leaf is a very inexpensive alternative, however aluminum is much more difficult to work than gold or silver and results in less than optimal results when employing traditional metal leafing technique, and so palladium leaf is considered the best substitute despite its considerable cost. Platinum leaf may be used to the same effect as palladium leaf with similar working properties, but it is not as readily available in leaf form commercially.[32] [33]
Safety Finely divided palladium metal can be pyrophoric. The bulk material is quite inert.
See also • • • • •
Palladium coin Precious metal Palladium as an investment Platinum Periodic Table
8
Palladium
9
External links • • • •
WebElements.com – Palladium [34] Current and Historical Palladium Price [35] picture of 999.5 fine Palladium in the element collection from Heinrich Pniok Comprehensive Data on Palladium [37]
[36]
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] " Palladium (http:/ / www. ipa-news. com/ pgm/ index. htm)". International Platinum Group Metals Association. . [3] W. P. Griffith (2003). " Rhodium and Palladium - Events Surrounding Its Discovery (http:/ / www. platinummetalsreview. com/ dynamic/ article/ view/ 47-4-175-183)". Platinum Metals Review 47 (4): 175–183. . [4] Wollaston, W. H. (1804). "On a New Metal, Found in Crude Platina". Philosophical Transactions of the Royal Society of London 94: 419–430. doi: 10.1098/rstl.1804.0019 (http:/ / dx. doi. org/ 10. 1098/ rstl. 1804. 0019). [5] C. R. Hammond. The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304857. [6] Garrett, Christine E.; Prasad, Kapa (2004). "The Art of Meeting Palladium Specifications in Active Pharmaceutical Ingredients Produced by Pd-Catalyzed Reactions". Advanced Synthesis & Catalysis 346 (8): 889–900. doi: 10.1002/adsc.200404071 (http:/ / dx. doi. org/ 10. 1002/ adsc. 200404071). [7] Alan Williamson. " Russian PGM Stocks (http:/ / www. lbma. org. uk/ conf2003/ 5d. williamson LBMAConf2003. pdf)". The LBMA Precious Metals Conference 2003. The London Bullion Market Association. . [8] " Historical Palladium Charts and Data (http:/ / www. kitco. com/ charts/ historicalpalladium. html)". Kitco. . Retrieved 2007-08-09. [9] " Ford fears first loss in a decade (http:/ / news. bbc. co. uk/ 1/ hi/ business/ 1763406. stm)". BBC News. 2002-01-16. . Retrieved 2008-09-19. [10] " Platinum-Group Metals (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ platinum/ platimcs07. pdf)". Mineral Commodity Summaries. United States Geological Survey. January 2007. . [11] J. Kielhorn, C. Melber, D. Keller, I. Mangelsdorf (2002). "Palladium – A review of exposure and effects to human health". International Journal of Hygiene and Environmental Health 205 (6): 417. doi: 10.1078/1438-4639-00180 (http:/ / dx. doi. org/ 10. 1078/ 1438-4639-00180). [12] Gray, Theodore. " 46 Palladium (http:/ / www. theodoregray. com/ periodictabledisplay/ Elements/ 046/ index. s9. html)". Element Displays. . Retrieved 2007-10-14. [13] W. R. Kelly, G. J. Wasserburg, (1978). "Evidence for the existence of 107Pd in the early solar system". Geophysical Research Letters 5: 1079–1082. doi: 10.1098/rsta.2001.0893 (http:/ / dx. doi. org/ 10. 1098/ rsta. 2001. 0893). [14] J. H. Chen, G. J. Wasserburg (1990). "The isotopic composition of Ag in meteorites and the presence of 107Pd in protoplanets". Geochimica et Cosmochimica Acta 54 (6): 1729–1743. doi: 10.1016/0016-7037(90)90404-9 (http:/ / dx. doi. org/ 10. 1016/ 0016-7037(90)90404-9). [15] Gordon K. Anderson, Minren Lin (1990). "Bis(Benzonitrile)Dichloro Complexes of Palladium and Platinum". Inorganic Syntheses 28: 60–63. doi: 10.1002/9780470132593.ch13 (http:/ / dx. doi. org/ 10. 1002/ 9780470132593. ch13). [16] " Palladium (http:/ / www. unctad. org/ infocomm/ anglais/ palladium/ uses. htm)". United Nations Conference on Trade and Development. . Retrieved 2007-02-05. [17] Roy Rushforth (2004). " Palladium in Restorative Dentistry: Superior Physical Properties make Palladium an Ideal Dental Metal (http:/ / www. platinummetalsreview. com/ dynamic/ article/ view/ 48-1-030-031)". Platinum Metals Review 48 (1). . [18] Hesse, Rayner W. (2007). " palladium (http:/ / books. google. de/ books?id=DIWEi5Hg93gC& pg=PA146)". Jewelrymaking through history: an encyclopedia. Greenwood Publishing Group. pp. 146. . [19] Toff, Nancy (1996). The flute book: a complete guide for students and performers (http:/ / books. google. de/ books?id=pCSanDD4CtsC& pg=PA20). Oxford University Press. pp. 20. ISBN 9780195105025. . [20] Dennis Zogbi (February 3, 2003). " Shifting Supply and Demand for Palladium in MLCCs (http:/ / www. ttiinc. com/ object/ ME_Zogbi_20030203. html)". TTI, Inc.. . [21] David Jollie (2007). " Platinum 2007 (http:/ / www. platinum. matthey. com/ uploaded_files/ 2007/ 07_palladium. pdf)" (PDF). Johnson Matthey. .
Palladium [22] Jiro Tsuji (2004). Palladium reagents and catalysts: new perspectives for the 21st century (http:/ / books. google. com/ books?id=RDT0OUdlj0MC& pg=PA90). John Wiley and Sons. p. 90. ISBN 0470850329. . [23] Chemical & Engineering News Vol. 86 No. 35, 1 Sept. 2008, "Palladium's Hidden Talent", p. 53 [24] W. Grochala, P. P. Edwards (2004). "Thermal Decomposition of the Non-Interstitial Hydrides for the Storage and Production of Hydrogen". Chem. Rev. 104 (3): 1283–1316. doi: 10.1021/cr030691s (http:/ / dx. doi. org/ 10. 1021/ cr030691s). [25] Emil Raymond Riegel, James A. Kent (2007). Kent and Riegel's Handbook of Industrial Chemistry and Biotechnology (http:/ / books. google. com/ books?id=zPZWYerB3SYC& pg=PA1037). Springer. p. 1037. ISBN 0387278427. . [26] " Cross-reactivity between nickel and palladium demonstrated by systemic administration of nickel. (http:/ / www. ncbi. nlm. nih. gov/ pubmed/ 15982224)". PubMed. . Retrieved 2009-06-06. [27] " What Is Palladium? (http:/ / www. jewelry. com/ palladium-what-is. shtml)". Jewelry.com. November 3, 2008. . [28] " Daily Metal Prices: September 2001 (http:/ / www. platinum. matthey. com/ prices/ September2001. php)". Johnson Matthey. . [29] Holmes, E. (February 13, 2007). "Palladium, Platinum's Cheaper Sister, Makes a Bid for Love". Wall Street Journal (Eastern edition). pp. B.1. [30] " Stillwater Mining Up on Jewelry Venture (http:/ / biz. yahoo. com/ ap/ 080311/ stillwater_mining_mover. html)". Yahoo Finance. . [31] Mike Ware (2005). "Book Review of : Photography in Platinum and Palladium". Platinum Metals Review 49 (4): 190–195. doi: 10.1595/147106705X70291 (http:/ / dx. doi. org/ 10. 1595/ 147106705X70291). [32] Margaret Morgan (2007). The Bible of Illuminated Letters. Barron's Educational Series. p. 50. ISBN 978-0764158209. [33] " Palladium Leaf (http:/ / www. theodoregray. com/ PeriodicTable/ Samples/ 046. 6/ index. s12. html)". Theodore Gray. . [34] [35] [36] [37]
http:/ / www. webelements. com/ webelements/ elements/ text/ Pd/ index. html http:/ / www. kitco. com/ charts/ livepalladium. html http:/ / www. pse-mendelejew. de/ bilder/ pd. jpg http:/ / www. mrteverett. com/ Chemistry/ pdictable/ q_elements. asp?Symbol=Pd
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Article Sources and Contributors
Article Sources and Contributors Palladium Source: http://en.wikipedia.org/w/index.php?oldid=307294061 Contributors: 56, AFewGoodMen, Achim1999, Adashiel, Ahoerstemeier, Alfio, AlimanRuna, Anclation, Andrewa, Anittas, Anwar saadat, Archimerged, Arteitle, Arthena, Astavats, AxelBoldt, Beetstra, Bigturtle, Blanchardb, BlueEarth, Borislav Dopudja, Braindigitalis, Brandonrush, Brian Huffman, Bryan Derksen, Bryancpark, Bung hole 77, Buttercup44, CYD, Carnildo, Chrisacip1419, Clement Cherlin, Conversion script, CopaceticThought, Coronos, Dan Allen, Darrien, Daru718, Davespice, David Latapie, David Richtman, Dcoetzee, Dead3y3, Deli nk, Delta G, Dennis Brown, Donarreiskoffer, Doulos Christos, Dr. Latex, Drphilharmonic, Drrocket, Dummies102, Dvorak729, Dwmyers, EPO, EarthPerson, Edgar181, Ekrumme, El, El C, Element16, Eloil, Emperorbma, Enviroboy, EoGuy, Epastore, Faradayplank, Fedallah, Felix Folio Secundus, Femto, Fg2, Flapdragon, Fromos, Froth, Gman124, Godlord2, Goldnanoparticles, Gracenotes, Grendelkhan, Gtrmp, Gurch, Hak-kâ-ngìn, Hallpriest9, Hardrockmine, Helge Skjeveland, Hobartimus, Hu12, IanOsgood, Icairns, Indefatigable, Infinoid, InvertRect, InvictaHOG, J.delanoy, JWB, JaGa, Jaffachief, Jaraalbe, Jason0416, Joanjoc, John, JohnCD, Jorge Stolfi, Jose77, JoshuaKuo, Jurj, Jwiki123, Kaosme, Karl-Henner, Kaverin, Kelson, King Lopez, Korath, Kwamikagami, LA2, Lbasse, Leafyplant, LeaveSleaves, LuigiManiac, MER-C, Magicbronson, Marhault, Materialscientist, Mav, McTrixie, Mdf, Melchoir, Mgimpel, Michael Hardy, Michael IFA, Michael L. Hall, Michal il, Minesweeper, Montrealais, Mrwordsmith, Natalie Erin, Nergaal, Neutrality, Nickstuckert, Night Gyr, Nihiltres, Nirvana2013, Nurg, One, Ourai, Ozzfinds, Pcarbonn, Peyre, Phe, Phileas, PierreAbbat, Plantsurfer, PlatinumX, Poolkris, Pras, Preciousmetalboy, Psyche825, RTC, Raul654, Raymondwinn, Remember, Renato Caniatti, Rich Farmbrough, Rifleman 82, Ringman17, Rjwilmsi, RobertG, Roberta F., Rojomoke, Romanm, Rootology, Rossman7000, Rossnorman, RupertOlson, SGGH, Saperaud, Savant13, Schneelocke, Sengkang, Shaddack, Shafei, Shatha, Sheitan, Shirulashem, Sillybilly, Silverbach, Silverbackmarlin, Sl, Sleigh, Slordak, Snurks, Sonett72, Starelda, Stefan da, Stephenb, Stifynsemons, Stone, Suisui, Sunborn, Sunderland06, Tagishsimon, Technicaltechy, Tetracube, Thegeneralguy, Time3000, Timeastor, Titus III, Toliver djs, UberScienceNerd, Ulric1313, Usien6, VASANTH S.N., Versus22, Vicarious, Vlad4599, Vsmith, Walkerma, Wayiran, Wereon, WikiPediaAid, Wizard191, WmRowan, Yakityak, Yinon, Yyy, Z.E.R.O., Zfr, 305 anonymous edits
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11
Silver
1
Silver 47
palladium ← silver → cadmium
Cu ↑
Ag ↓
Au Periodic Table - Extended Periodic Table
General Name, Symbol, Number
silver, Ag, 47
Element category
transition metals
Group, Period, Block
11, 5, d
Appearance
lustrous white metal
Standard atomic weight
107.8682(2) g·mol
Electron configuration
[Kr] 4d
−1
10
Electrons per shell
1
5s
2, 8, 18, 18, 1 Physical properties
Color
silver
Phase
solid
Density (near r.t.)
10.49 g·cm−3
Liquid density at m.p.
9.320 g·cm−3
Melting point
1234.93 K (961.78 °C, 1763.2 °F)
Boiling point
2435 K (2162 °C, 3924 °F)
Heat of fusion
11.28 kJ·mol−1
Heat of vaporization
250.58 kJ·mol−1
Specific heat capacity
(25 °C) 25.350 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1283
1413
1575
1782
2055
2433
Atomic properties
Silver
2
Crystal structure
face-centered cubic
Oxidation states
1, 2, 3 (amphoteric oxide)
Electronegativity Ionization energies
1.93 (Pauling scale) 1st: 731.0 kJ/mol 2nd: 2070 kJ/mol 3rd: 3361 kJ/mol
Atomic radius
144 pm
Covalent radius
145±5 pm
Van der Waals radius
172 pm Miscellaneous
Magnetic ordering
diamagnetic
Electrical resistivity
(20 °C) 15.87 n Ω·m
Thermal conductivity
(300 K) 429 W·m
Thermal diffusivity
(300 K) 174 mm²/s
Thermal expansion
(25 °C) 18.9 µm·m
Speed of sound (thin rod)
(r.t.) 2680 m·s
Young's modulus
83 GPa
Shear modulus
30 GPa
Bulk modulus
100 GPa
−1
−1
−1
Poisson ratio
0.37
Mohs hardness
2.5
Vickers hardness
251 MPa
Brinell hardness
24.5 MPa
CAS registry number
·K
7440-22-4 Most-stable isotopes
−1
−1
·K
Silver
3
Main article: Isotopes of silver iso 105
Ag
NA syn
half-life 41.2 d
DM ε
-
γ
106m
Ag
syn
8.28 d
ε
Ag
108m
Ag
51.839% syn
Ag
111
Ag
48.161% syn
0.511, 0.717, 1.045, 0.450
DP 105
Pd
-
106
Pd
-
107
Ag is stable with 60 neutron
418 y
ε
-
108
IT
0.109
108
γ
109
0.344, 0.280, 0.644, 0.443 -
γ
107
DE (MeV)
0.433, 0.614, 0.722
Pd Ag
-
109
Ag is stable with 62 neutron
7.45 d
β
1.036, 0.694
γ
0.342
−
111
Cd
-
References
Silver is a chemical element with the chemical symbol Ag (Latin: argentum, from the Ancient Greek: ἀργήεντος - argēentos, gen. of ἀργήεις - argēeis, "white, shining" ) and atomic number 47. A soft, white, lustrous transition metal, it has the highest electrical conductivity of any element and the highest thermal conductivity of any metal. The metal occurs naturally in its pure, free form (native silver), as an alloy with gold (electrum) and other metals, and in minerals such as argentite and chlorargyrite. Most silver is produced as a by-product of copper, gold, lead, and zinc refining. Silver has been known since ancient times and has long been valued as a precious metal, used to make ornaments, jewelry, high-value tableware, utensils (hence the term silverware), and currency coins. Today, silver metal is used in electrical contacts and conductors, in mirrors and in catalysis of chemical reactions. Its compounds are used in photographic film and dilute solutions of silver nitrate and other silver compounds are used as disinfectants. Although the antimicrobial uses of silver have largely been supplanted by the use of antibiotics, further research into its clinical potential is in progress.
Silver
4
Characteristics
Silver 1000oz bullion bar
Silver is a very ductile and malleable (slightly harder than gold) monovalent coinage metal with a brilliant white metallic luster that can take a high degree of polish. It has the highest electrical conductivity of all metals, even higher than copper, but its greater cost and tarnishability have prevented it from being widely used in place of copper for electrical purposes, though 13540 tons were used in the electromagnets used for enriching uranium during World War II (mainly because of the wartime shortage of copper).[1] [2] Another notable exception is in high-end audio cables.[3]
Among metals, pure silver has the highest thermal conductivity[4] (the non-metal diamond and superfluid helium II are higher), the whitest color, and the highest optical reflectivity[5] (although aluminium slightly outdoes it in parts of the visible spectrum, and it is a poor reflector of ultraviolet light). Silver also has the lowest contact resistance of any metal. Silver halides are photosensitive and are remarkable for their ability a silver crystal, electrolytic refined to record a latent image that can later be developed with visible dendritic structures. chemically. Silver is stable in pure air and water, but tarnishes when it is exposed to air or water containing ozone or hydrogen sulfide to form a black layer of silver sulfide which can be cleaned off with dilute hydrochloric acid.[6] The most common oxidation state of silver is +1 (for example, silver nitrate: AgNO3); in addition, +2 compounds (for example, silver(II) fluoride: AgF2) and +3 compounds (for example, potassium tetrafluoroargentate: K[AgF4]) are known.
Isotopes Naturally occurring silver is composed of two stable isotopes, 107Ag and 109Ag, with 107Ag being the most abundant (51.839% natural abundance). Silver's standard atomic mass is 107.8682(2) u. Twenty-eight radioisotopes have been characterized, the most stable being 105 Ag with a half-life of 41.29 days, 111Ag with a half-life of 7.45 days, and 112Ag with a half-life of 3.13 hours. This element has numerous meta states, the most stable being 108m Ag (t* 418 years), 110mAg (t* 249.79 days) and 106mAg (t* 8.28 days). All of the remaining radioactive isotopes have half-lives that are less than an hour, and the majority of these have half-lives that are less than 3 minutes. Isotopes of silver range in atomic weight from 93.943 u (94Ag) to 123.929 u (124Ag). The primary decay mode before the most abundant stable isotope, 107Ag, is electron capture and the primary mode after is beta decay. The primary decay products before 107Ag are palladium (element 46) isotopes, and the primary products after are cadmium (element 48) isotopes. The palladium isotope 107Pd decays by beta emission to 107Ag with a half-life of 6.5 million years. Iron meteorites are the only objects with a high-enough palladium-to-silver ratio to yield measurable variations in 107Ag abundance. Radiogenic 107Ag was first discovered in
Silver
5
the Santa Clara meteorite in 1978.[7] The discoverers suggest that the coalescence and differentiation of iron-cored small planets may have occurred 10 million years after a nucleosynthetic event. 107Pd–107Ag correlations observed in bodies that have clearly been melted since the accretion of the solar system must reflect the presence of unstable [8] nuclides in the early solar system.
Silver compounds Silver metal dissolves readily in nitric acid (HNO3) to produce silver nitrate (AgNO3), a transparent crystalline solid that is photosensitive and readily soluble in water. Silver nitrate is used as the starting point for the synthesis of many other silver compounds, as an antiseptic, and as a yellow stain for glass in stained glass. Silver metal does not react with sulfuric acid, which is used in jewellery-making to clean and remove copper oxide firescale from silver articles after silver soldering or annealing. However, silver reacts readily with sulfur or hydrogen sulfide H2S to produce silver sulfide, a dark-coloured compound familiar as the tarnish on silver coins and other objects. Silver sulfide also forms silver whiskers when silver electrical contacts are used in an atmosphere rich in hydrogen sulfide. Silver bromide is a yellow, low hardness salt. Silver chloride (AgCl) is precipitated from solutions of silver nitrate in the presence of chloride ions, and the other silver halides used in the manufacture of photographic emulsions are made in the same way using bromide or iodide salts. Silver chloride is used in glass electrodes for pH testing and potentiometric measurement, and as a transparent cement for glass. Silver iodide has been used in attempts to seed clouds to produce rain.[6]
Cessna 210 equipped with a silver iodide generator for cloud seeding
Silver oxide (Ag O) produced when silver nitrate 2
solutions are treated with a base, is used as a positive electrode (cathode) in watch (battery) batteries. Silver carbonate (Ag2CO3) is precipitated when silver nitrate is treated with sodium carbonate (Na2CO3).[9] Silver fulminate (AgONC), a powerful, touch-sensitive explosive used in percussion caps, is made by reaction of silver metal with nitric acid in the presence of ethanol (C2H5OH). Another dangerously explosive silver compound is silver azide (AgN3), formed by reaction of silver nitrate with sodium azide (NaN3).[10] Latent images formed in silver halide crystals are developed by treatment with alkaline solutions of reducing agents such as hydroquinone, metol (4-(methylamino)phenol sulfate) or ascorbate which reduce the exposed halide to silver metal. Alkaline solutions of silver nitrate can be reduced to silver metal by reducing sugars such as glucose, and this reaction is used to silver glass mirrors and the interior of glass Christmas ornaments. Silver halides are soluble in solutions of sodium thiosulfate (Na2S2O3) which is used as a photographic fixer, to remove excess silver halide from photographic emulsions after image development.[9] Silver metal is attacked by strong oxidizers such as potassium permanganate (KMnO4) and potassium dichromate (K2Cr2O7), and in the presence of potassium bromide (KBr), these compounds are used in photography to bleach silver images, converting them to silver
Silver halides that can either be fixed with thiosulfate or re-developed to intensify the original image. Silver forms cyanide complexes (silver cyanide) that are soluble in water in the presence of an excess of cyanide ions. Silver cyanide solutions are used in electroplating of silver.[9]
Applications Precious metal A major use of silver is as a precious metal, and it has long been used for making high-value objects reflecting the wealth and status of the owner. For example the idiom "[to be] born with a Silver Spoon in one's mouth". Currency Silver, in the form of electrum (a gold-silver alloy), was coined to produce money in around 700 BCE by the Lydians. Later, silver was refined and coined in its pure Goddess Minerva on a Roman silver form. Many nations used silver as the basic unit of plate, 1st century BCE monetary value. The words for "silver" and "money" are the same in at least 14 languages. In the modern world, silver bullion has the ISO currency code XAG. The name of the United Kingdom monetary unit "pound" reflects the fact that it originally represented the value of one troy pound of sterling silver. In the 1800s, many nations, such as the United States and Great Britain, switched from silver to a gold standard of monetary value, then in the 20th century to fiat currency. Jewellery and silverware Jewellery and silverware are traditionally made from sterling silver (standard silver), an alloy of 92.5% silver with 7.5% copper. In the United States, only an alloy consisting of at least 92.5% fine silver can be marketed as "silver". Sterling silver is harder than pure silver, and has a lower melting point (893 °C) than either pure silver or pure copper.[6] Britannia silver is an alternative hallmark-quality standard containing 95.8% silver, often used to make silver tableware and wrought plate. With the addition of germanium, the patented modified alloy Argentium Sterling Silver is formed, with improved properties including resistance to firescale. Sterling silver jewelry is often plated with a thin coat of .999 fine silver to give the item a shiny finish. This process is called "flashing". Silver jewelry can also be plated with rhodium (for a bright, shiny look) or gold. Silver is a constituent of almost all colored carat gold alloys and carat gold solders, giving the alloys paler colour and greater hardness.[11] White 9 carat gold contains 62.5% silver and 37.5% gold, while 22 carat gold contains up to 8.4% silver or 8.4% copper.[11] Historically the training and guild organization of goldsmiths included silversmiths as well, and the two crafts remain largely overlapping. Unlike blacksmiths, silversmiths do not shape the metal while it is red-hot but instead, work it at room temperature with gentle and carefully placed hammerblows. The essence of silversmithing is to take a flat piece of metal and by means of different hammers, stakes and other simple tools, to transform it into a
6
Silver useful object.[12] While silversmiths specialize in, and principally work, silver, they also work with other metals such as gold, copper, steel, and brass. They make jewellery, silverware, armour, vases, and other artistic items. Because silver is such a malleable metal, silversmiths have a large range of choices with how they prefer to work the metal. Historically, silversmiths are mostly referred to as goldsmiths, which was usually the same guild. In the western Canadian silversmith tradition, guilds do not exist; however, mentoring through colleagues becomes a method of professional learning within a community of craftspeople.[13] Silver is cheaper than gold, though still valuable, and so is very popular with jewellers who are just starting out and cannot afford to make pieces in gold, or as a practicing material for goldsmith apprentices. Silver has also become very fashionable, and is used frequently in more artistic jewellery pieces. Traditionally silversmiths mostly made "silverware" (cutlery, table flatware, bowls, candlesticks and such). Only in more recent times has silversmithing become mainly work in jewellery, as much less solid silver tableware is now handmade. Silver is used in medals, denoting second place. Some high-end musical instruments are made from sterling silver, such as the flute.
Dentistry Silver can be alloyed with mercury, tin and other metals at room temperature to make amalgams that are widely used for dental fillings. To make dental amalgam, a mixture of powdered silver and other metals is mixed with mercury to make a stiff paste that can be adapted to the shape of a cavity. The dental amalgam achieves initial hardness within minutes but sets hard in a few hours.
Photography and electronics Photography used 30.98% of the silver consumed in 1998 in the form of silver nitrate and silver halides. In 2001, 23.47% was used for photography, while 20.03% was used in jewelry, 38.51% for industrial uses, and only 3.5% for coins and medals. The use of silver in photography has rapidly declined, due to the lower demand for consumer colour film from the advent of digital technology, since in 2007 of the 894.5 million ounces of silver in supply, just 128.3 million ounces (14.3%) were consumed by the photographic sector, and the total amount of silver consumed in 2007 by the photographic sector compared to 1998 is just 50%.[14] Some electrical and electronic products use silver for its superior conductivity, even when tarnished. For example, printed circuits are made using silver paints,[6] and computer keyboards use silver electrical contacts. Some high-end audio hardware (DACs, preamplifiers, etc.) are fully silver-wired, which is believed to cause the least loss of quality in the signal. Silver cadmium oxide is used in high voltage contacts because it can withstand arcing. During World War II the short supply of copper brought about the government's use of silver from the Treasury vaults for conductors at Oak Ridge National Laboratory. (After the war ended the silver was returned to the vaults.)[15]
7
Silver
Solder and brazing Silver is used to make solder and brazing alloys, electrical contacts, and high-capacity silver-zinc and silver-cadmium batteries. Silver in a thin layer on top of a bearing material can provide a significant increase in galling resistance and reduce wear under heavy load, particularly against steel.
Mirrors and optics Mirrors which need superior reflectivity for visible light are made with silver as the reflecting material in a process called silvering, though common mirrors are backed with aluminium. Using a process called sputtering, silver (and sometimes gold) can be applied to glass at various thicknesses, allowing different amounts of light to penetrate. Silver is usually reserved for coatings of specialized optics, and the silvering most often seen in architectural glass and tinted windows on vehicles is produced by sputtered aluminium, which is cheaper and less susceptible to tarnishing and corrosion.[16]
Nuclear reactors Because silver readily absorbs free neutrons, it is commonly used to make control rods that regulate the fission chain reaction in pressurized water nuclear reactors, generally in the form of an alloy containing 80% silver, 15% indium, and 5% cadmium.
Catalyst Silver's catalytic properties make it ideal for use as a catalyst in oxidation reactions, for example, the production of formaldehyde from methanol and air by means of silver screens or crystallites containing a minimum 99.95 weight-percent silver. Silver (upon some suitable support) is probably the only catalyst available today to convert ethylene to ethylene oxide (later hydrolyzed to ethylene glycol, used for making polyesters)—a very important industrial reaction. Oxygen dissolves in silver relatively easily compared to other gases present in air. Attempts have been made to construct silver membranes of only a few monolayers thickness. Such a membrane could be used to filter pure oxygen from air and water.
8
Silver
9
Medicine Silver ions and silver compounds show a toxic effect on some bacteria, viruses, algae and fungi, typical for heavy metals like lead or mercury, but without the high toxicity to humans that are normally associated with these other metals. Its germicidal effects kill many microbial organisms in vitro, but testing and standardization of silver products is difficult.[17] Hippocrates, the father of modern medicine, wrote that silver had beneficial healing and anti-disease properties, and the Phoenicians used to store water, wine, and vinegar in silver bottles to prevent spoiling. In the early 1900s people would put silver dollars in milk bottles to prolong the milk's freshness.[18] Its germicidal effects increased its value in utensils and as jewellery. The exact process of silver's germicidal effect is still not well understood, although theories exist. One of these is the oligodynamic effect, which explains the effect on microorganisms but would not explain antiviral effects.
Handforged silver wine goblets. Usage of silverware was encouraged by the antibiotic action of silver
Silver compounds were used to prevent infection in World War I before the advent of antibiotics. Silver nitrate solution was a standard of care but was largely replaced by silver sulfadiazine cream (SSD Cream),[19] which was generally the "standard of care" for the antibacterial and antibiotic treatment of serious burns until the late 1990s.[20] Now, other options, such as silver-coated dressings (activated silver dressings), are used in addition to SSD cream. However, the evidence for the effectiveness of such silver-treated dressings is mixed and although the evidence is promising it is marred by the poor quality of the trials used to assess these products.[21] Consequently a major systematic review by the Cochrane Collaboration found insufficient evidence to recommend the use of silver-treated dressings to treat infected wounds.[21] The widespread use of silver went out of fashion with the development of modern antibiotics. However, recently there has been renewed interest in silver as a broad-spectrum antimicrobial agent. In particular, silver is being used with alginate, a naturally occurring biopolymer derived from seaweed, in a range of products designed to prevent infections as part of wound management procedures, particularly applicable to burn victims.[22] In 2007, AGC Flat Glass Europe introduced the first antibacterial glass to fight hospital-caught infection: it is covered with a thin layer of silver.[23] In addition, the U.S. Food and Drug Administration (FDA) has recently approved an endotracheal breathing tube with a fine coat of silver for use in mechanical ventilation, after studies found it reduced the risk of ventilator-associated pneumonia.[24] It has long been known that antibacterial action of silver is enhanced by the presence of an electric field. Applying a few volts of electricity across silver electrodes drastically enhances the rate that bacteria in solution are killed. It was found recently that the antibacterial action of silver electrodes is greatly improved if the electrodes are covered with silver nanorods.[25]
Silver Medication Today, various kinds of silver compounds, or devices to make solutions or colloids containing silver, are sold as remedies for a wide variety of diseases. Although most colloidal silver preparations are harmless, some people using these home-made solutions excessively have developed argyria over a period of months or years.[26] High doses of colloidal silver can result in coma, pleural edema, and hemolysis.[27] Silver is widely used in topical gels and impregnated into bandages because of its wide-spectrum antimicrobial activity. The anti-microbial properties of silver stem from the chemical properties of its ionized form, Ag+. This ion forms strong molecular bonds with other substances used by bacteria to respire, such as molecules containing sulfur, nitrogen, and oxygen.[28] Once the Ag+ ion complexes with these molecules, they are rendered unusable by the bacteria, depriving it of necessary compounds and eventually leading to the bacteria's death.
Food Silver is now considered to be a bioaccumulating carcinogen and thus its use in food is not advised. In India and Pakistan, foods, especially sweets, can be found decorated with a thin layer of silver known as vark. Silver as a food additive is given the E number E174 and is classed as a food coloring. It is used solely for external decoration, such as on chocolate confectionery, in the covering of dragées and the decoration of sugar-coated flour confectionery. In Australia, it is banned as a food additive.
Clothing Silver inhibits the growth of bacteria and fungi. It keeps odor to a minimum and reduces the risk of bacterial and fungal infection. In clothing, the combination of silver and moisture movement (wicking) may help to reduce the harmful effects of prolonged use in active and humid conditions. Silver is used in clothing in two main forms: • A form in which silver ions are integrated into the polymer from which yarns are made (a form of nanotechnology) • A form in which the silver is coated onto the yarns. In both cases the silver prevents the growth of a broad spectrum of bacteria and fungi. Recorded use of silver to prevent infection dates to ancient Greece and Rome, it was rediscovered in the Middle Ages, where it was used for several purposes, such as to disinfect water and food during storage, and also for the treatment of burns and wounds as wound dressing. In the 19th century, sailors on long ocean voyages would put silver coins in barrels of water and wine to keep the liquid pure. Pioneers in America used the same idea as they made their journey from coast to coast. Silver solutions were approved in the 1920s by the US Food and Drug Administration for use as antibacterial agents. Today, wound dressings containing silver are well established for clinical wound care and have recently been introduced in consumer products such as sticking plasters.[29]
10
Silver
11
History
The symbol for the Moon has been used since ancient times to represent silver.
The word "silver" appears in Anglo-Saxon in various spellings such as seolfor and siolfor. A similar form is seen throughout the Teutonic languages (compare Old High German silabar and silbir). The chemical symbol Ag is from the Latin for "silver", argentum (compare Greek άργυρος, árgyros), from the Indo-European root *arg- meaning "white" or "shining". Silver has been known since ancient times. It is mentioned in the book of Genesis, and slag heaps found in Asia Minor and on the islands of the Aegean Sea indicate that silver was being separated from lead as early as the 4th millennium BC using surface mining.[6] In the Gospels, Jesus' disciple Judas Iscariot is infamous for having taken a bribe of thirty coins of silver from religious leaders in Jerusalem to turn Jesus Christ over
to the Romans. Set aside certain circumstances, Islam permits Muslim men to wear silver jewelry. Muhammad himself wore a silver signet ring.
Occurrence and extraction Silver is found in native form, alloyed with gold or combined with sulfur, arsenic, antimony or chlorine in ores such as argentite (Ag2S), horn silver (AgCl), and pyrargyrite (Ag3SbS3). The principal sources of silver are the ores of copper, copper-nickel, lead, and lead-zinc obtained from Peru, Mexico, China, Australia, Chile, Poland and Serbia.[6] Peru and Mexico have been mining silver since 1546 and are still major world producers. Top silver-producing mines are Proaño / Fresnillo (Mexico), Cannington (Queensland, Australia), Dukat (Russia), Uchucchacua (Peru) and Greens Creek mine (Alaska).[30] The metal can also be produced during the electrolytic refining of copper and by the application of the Parkes process on lead metal obtained from lead ores that contain small amounts of silver. Commercial-grade fine silver is at least 99.9% pure silver, and purities greater than 99.999% are available. In 2007, Peru was the world's top producer of silver, closely followed by Mexico, according to the British Geological Survey.
Silver ore with Lincoln cent for scale
Time trend of silver production
Silver
12
Price As of October 2008 silver is about 1/75th the price of gold by mass.[31] Silver once traded at 1/6th to 1/12th the price of gold, prior to the Age of Discovery and the discovery of great silver deposits in the Americas, including Peru, Mexico and the United States, such as the vast Comstock Silver output in 2005 Lode in Virginia City, Nevada, USA. This then resulted in the debate over cheap Free Silver to benefit the agricultural sector, which was among the most prolonged and difficult in that country's history and dominated public discourse during the latter decades of the nineteenth century. Over the last 100 years the price of silver and the gold/silver price ratio have fluctuated greatly due to competing industrial and store-of-value demands. In 1980 the silver price rose to an all-time high of US$49.45 per troy ounce. By December 2001 the price had dropped to US$4.15 per ounce, and in May 2006 it had risen back as high as US$15.21 per ounce. In March 2008 silver reached US$21.34 per ounce.[32] The price of silver is important in Judaic Law. The lowest fiscal amount that a Jewish court, or Beth Din, can convene to adjudicate a case over is a shova pruta (value of a Babylonian pruta coin). This is fixed at 1/8 of a gram of pure, unrefined silver, at market price.
Precautions Silver plays no known natural biological role in humans, and possible health effects of silver are a subject of dispute. Silver itself is not toxic but most silver salts are, and some may be carcinogenic. Silver and compounds containing silver (like colloidal silver) can be absorbed into the circulatory system and become deposited in various body tissues leading to a condition called argyria which results in a blue-grayish pigmentation of the skin, eyes, and mucous membranes. Although this condition does not otherwise harm a person's health, it is disfiguring and usually permanent. Argyria is rare, and mild forms are sometimes mistaken for cyanosis.[6]
See also • List of Silver Compounds • Silverpoint drawing
External links • WebElements.com – Silver [33] • Society of American Silversmiths
[34]
• Silver Statistics and Information [35], USGS publications on the worldwide production of silver • The Silver Institute [36] A silver industry website
Silver • A collection of silver items [37] Samples of silver • Transport, Fate and Effects of Silver in the Environment [38] • Picture in the Element collection from Heinrich Pniok [39]
References [1] Nichols, Kenneth D. (1987). The Road to Trinity. Morrow, New York: Morrow. p. 42. ISBN 068806910X. [2] " Eastman at Oak Ridge - Dr. Howard Young (http:/ / www. tnengineering. net/ AICHE/ eastman-oakridge-young. htm)". . Retrieved 2009-06-06. [3] Oman, H. (1992). "Not invented here? Check your history". Aerospace and Electronic Systems Magazine 7 (1): 51–53. doi: 10.1109/62.127132 (http:/ / dx. doi. org/ 10. 1109/ 62. 127132). [4] " WebElements Periodic Table of the Elements | Silver | Essential information (http:/ / www. webelements. com/ silver/ )". Webelements.com. . Retrieved 2009-04-05. [5] Edwards, H.W. & Petersen, R.P. (1936). "Reflectivity of evaporated silver films". Phys. Rev. 9: 871. [6] C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304814. [7] Kelly, William R. (1978). "Evidence for the existence of 107Pd in the early solar system". Geophysical Research Letters 5: 1079. doi: 10.1029/GL005i012p01079 (http:/ / dx. doi. org/ 10. 1029/ GL005i012p01079). [8] Russell, Sara S.; Gounelle, Matthieu; Hutchison, Robert (2001 pages = 1991–2004). " Origin of Short-Lived Radionuclides (http:/ / www. jstor. org/ stable/ 3066270)". Philosophical Transactions: Mathematical, Physical and Engineering Sciences 359 (1787). doi:10.2307/3066270 (inactive 2009-05-30) . . [9] Hans I. Bjelkhagen (1995). Silver-halide recording materials: for holography and their processing. Springer. pp. 156-166. ISBN 3540586199. [10] Rudolf Meyer, Josef Köhler, Axel Homburg (2007). Explosives (http:/ / books. google. com/ books?id=ATiYCfo1VcEC& pg=PA284& dq="silver+ + fulminate"& lr=& num=100& as_brr=3). Wiley-VCH. p. 284. ISBN 3527316566. . [11] " Gold Jewellery Alloys > Utilise Gold. Scientific, industrial and medical applications, products ,suppliers from the World Gold Council (http:/ / www. utilisegold. com/ jewellery_technology/ colours/ colour_alloys/ )". Utilisegold.com. 2000-01-20. . Retrieved 2009-04-05. [12] " Chambers Search Chambers (http:/ / www. chambersharrap. co. uk/ chambers/ features/ chref/ chref. py/ main?query=Silversmith& title=21st& sourceid=Mozilla-search)". . Retrieved 2009-06-06. [13] Kelly McRae. " Trade Secrets (http:/ / westernhorseman. com/ index. php?option=com_content& task=view& id=861& Itemid=79)". Western Horseman Magazine. . Retrieved 2009-06-06. [14] " Silver Supply & Demand (http:/ / www. silverinstitute. org/ supply_demand. php)". The Silver Institute. . Retrieved 2009-05-05. [15] Isaac Asimov. Building Blocks of the Universe. [16] Ray N. Wilson (2004). Reflecting Telescope Optics: Basic design theory and its historical development (http:/ / books. google. com/ books?id=isH9fTnpc7YC& lpg=PA429& dq=silver optics& lr=& as_drrb_is=q& as_minm_is=0& as_miny_is=& as_maxm_is=0& as_maxy_is=& num=100& as_brr=3& pg=PA429). Springer. p. 422. ISBN 3540401067. . [17] Chopra I (April 2007). "The increasing use of silver-based products as antimicrobial agents: a useful development or a cause for concern?". The Journal of antimicrobial chemotherapy 59 (4): 587–90. doi: 10.1093/jac/dkm006 (http:/ / dx. doi. org/ 10. 1093/ jac/ dkm006). PMID 17307768. [18] " Antibacterial effects of silver (http:/ / www. saltlakemetals. com/ Silver_Antibacterial. htm)". . [19] Chang TW, Weinstein L (December 1975). " Prevention of herpes keratoconjunctivitis in rabbits by silver sulfadiazine (http:/ / aac. asm. org/ cgi/ pmidlookup?view=long& pmid=1211919)". Antimicrob. Agents Chemother. 8 (6): 677–8. PMID 1211919. PMC: 429446 (http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?tool=pmcentrez& artid=429446). . [20] Atiyeh BS, Costagliola M, Hayek SN, Dibo SA (March 2007). "Effect of silver on burn wound infection control and healing: review of the literature". Burns : journal of the International Society for Burn Injuries 33 (2): 139–48. doi: 10.1016/j.burns.2006.06.010 (http:/ / dx. doi. org/ 10. 1016/ j. burns. 2006. 06. 010). PMID 17137719. [21] Lo SF, Hayter M, Chang CJ, Hu WY, Lee LL (August 2008). "A systematic review of silver-releasing dressings in the management of infected chronic wounds". Journal of clinical nursing 17 (15): 1973–85. doi: 10.1111/j.1365-2702.2007.02264.x (http:/ / dx. doi. org/ 10. 1111/ j. 1365-2702. 2007. 02264. x). PMID 18705778. [22] Hermans MH (December 2006). "Silver-containing dressings and the need for evidence". The American journal of nursing 106 (12): 60–8; quiz 68–9. PMID 17133010.
13
Silver
14
[23] " AGC Flat Glass Europe launches world’s first antibacterial glass (http:/ / www. agc-flatglass. eu/ AGC+ Flat+ Glass+ Europe/ English/ Homepage/ News/ Press+ room/ Press-Detail-Page/ page. aspx/ 979?pressitemid=1031)". 2007-09-04. . [24] " FDA Clears Silver-Coated Breathing Tube For Marketing (http:/ / www. fda. gov/ bbs/ topics/ NEWS/ 2007/ NEW01741. html)". 2007-11-08. . Retrieved 2007-11-11. [25] O. Akhavan and E. Ghaderi (2009). " Enhancement of antibacterial properties of Ag nanorods by electric field (http:/ / www. iop. org/ EJ/ article/ 1468-6996/ 10/ 1/ 015003/ stam9_1_015003. pdf)". Sci. Technol. Adv. Mater. 10: 015003. doi: 10.1088/1468-6996/10/1/015003 (http:/ / dx. doi. org/ 10. 1088/ 1468-6996/ 10/ 1/ 015003). . [26] Fung MC, Bowen DL (1996). "Silver products for medical indications: risk-benefit assessment". Journal of toxicology. Clinical toxicology 34 (1): 119–26. doi: 10.3109/15563659609020246 (http:/ / dx. doi. org/ 10. 3109/ 15563659609020246). PMID 8632503. [27] Wadhera A, Fung M (2005). " Systemic argyria associated with ingestion of colloidal silver (http:/ / dermatology. cdlib. org/ 111/ case_reports/ argyria/ wadhera. html)". Dermatology online journal 11 (1): 12. PMID 15748553. . [28] Slawson RM, Van Dyke MI, Lee H, Trevors JT (January 1992). "Germanium and silver resistance, accumulation, and toxicity in microorganisms". Plasmid 27 (1): 72–9. doi: 10.1016/0147-619X(92)90008-X (http:/ / dx. doi. org/ 10. 1016/ 0147-619X(92)90008-X). PMID 1741462. [29] " Silver in Textiles and Clothing (http:/ / www. americanelements. com/ agnp. html)". Americanelements.com. . Retrieved 2009-04-05. [30] " Top silver producers (http:/ / www. infomine. com/ commodities/ silver. asp)". Infomine.com. . Retrieved 2009-04-05. [31] O’Connell, Rhona. " Gold:silver ratio will narrow, but base metals outlook ultra-gloomy – UBS (http:/ / www. mineweb. com/ mineweb/ view/ mineweb/ en/ page33?oid=71867& sn=Detail)". Mineweb.com. . Retrieved 2008-11-13. [32] " Silver Cash daily plot (http:/ / charts3. barchart. com/ chart. asp?sym=SIY0& data=A& jav=adv& vol=Y& evnt=adv& grid=Y& code=BSTK& org=stk& fix=)". Barchart.com. . [33] [34] [35] [36] [37] [38] [39]
http:/ / www. webelements. com/ webelements/ elements/ text/ Ag/ index. html http:/ / www. silversmithing. com http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ silver/ http:/ / www. silverinstitute. org http:/ / www. theodoregray. com/ PeriodicTable/ Elements/ 047/ index. html http:/ / digital. library. wisc. edu/ 1711. dl/ EcoNatRes. Argentum http:/ / www. pniok. de/ ag. htm
Article Sources and Contributors
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15
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16
Cadmium
1
Cadmium 48
silver ← cadmium → indium
Zn ↑
Cd ↓
Hg Periodic Table - Extended Periodic Table
General Name, Symbol, Number
cadmium, Cd, 48
Element category
transition metals
Group, Period, Block
12, 5, d
Appearance
silvery gray metallic
Standard atomic weight
112.411(8) g·mol
Electron configuration
[Kr] 5s 4d
−1
2
Electrons per shell
10
2, 8, 18, 18, 2 Physical properties
Phase
solid
Density (near r.t.)
8.65 g·cm
Liquid density at m.p.
7.996 g·cm−3
Melting point
594.22 K (321.07 °C, 609.93 °F)
Boiling point
1040 K (767 °C, 1413 °F)
Heat of fusion
6.21 kJ·mol−1
Heat of vaporization
99.87 kJ·mol−1
Specific heat capacity
(25 °C) 26.020 J·mol−1·K−1
−3
Vapor pressure P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
530
583
654
745
867
1040
Atomic properties Crystal structure
hexagonal
Cadmium
Oxidation states Electronegativity Ionization energies
2 2, 1 (mildly basic oxide) 1.69 (Pauling scale) 1st: 867.8 kJ/mol 2nd: 1631.4 kJ/mol 3rd: 3616 kJ/mol
Atomic radius
151 pm
Covalent radius
144±9 pm
Van der Waals radius
158 pm Miscellaneous
Magnetic ordering Electrical resistivity
[1]
diamagnetic
(22 °C) 72.7 nΩ·m
Thermal conductivity
(300 K) 96.6 W·m
Thermal expansion
(25 °C) 30.8 µm·m
Speed of sound (thin rod)
(20 °C) 2310 m/s
Young's modulus
50 GPa
Shear modulus
19 GPa
Bulk modulus
42 GPa
−1
−1
Poisson ratio
0.30
Mohs hardness
2.0
Brinell hardness CAS registry number
203 MPa 7440-43-9 Most-stable isotopes
−1
·K
−1
·K
Cadmium
3
Main article: Isotopes of cadmium iso 106
Cd
107
Cd
108
Cd
109
Cd
NA 1.25% syn 0.89% syn
half-life
-
106
6.5 h
ε
1.417
107
>6.7×1017 y
εε2ν
-
108
462.6 d
ε
0.214
109
−
0.316
113
β
−
0.580
113
IT
0.264
113
ββ2ν
-
114
β
1.446
115
ββ2ν
-
116
110
111
12.8%
111
112
24.13%
112
113
12.22%
Cd Cd
113m
Cd
114
Cd
115
Cd
116
Cd
syn
28.73% syn 7.49%
DP
εε2ν
12.49%
Cd
DE (MeV)
>9.5×1017 y
110
Cd
DM
Pd Ag Pd Ag
Cd is stable with 62 neutron Cd is stable with 63 neutron Cd is stable with 64 neutron 15
7.7×10
y
β
14.1 y
17
>9.3×10 53.46 h
19
2.9×10
y
−
y
In In Cd Sn In Sn
References
Cadmium (pronounced /ˈkædmiəm/) is a chemical element with the symbol Cd and atomic number 48. The soft, bluish-white transition metal is chemically similar to the two other metals in group 12, zinc and mercury. Similar to zinc it prefers oxidation state +2 in most of its compounds and similar to mercury it shows a low melting point for a transition metal. Cadmium is a relatively abundant element. Cadmium was discovered in 1817 by Friedrich Strohmeyer as an impurity in zinc carbonate. Cadmium occurs as a minor component in most zinc ores and therefore is a by-product of zinc production. Cadmium was for a long time used as pigment and for corrosion resistant plating on steel. Cadmium compounds were used to stabilize plastic. With the exception of its use in nickel-cadmium batteries, the use of cadmium is generally decreasing in all other applications. This decrease is due to the high toxicity and carcinogenicity of cadmium and the associated health and environmental concerns. Although cadmium is toxic, one enzyme, a carbonic anhydrase with a cadmium as reactive centre has been discovered.
Cadmium
4
Characteristics Cadmium is a soft, malleable, ductile, toxic, bluish-white bivalent metal. It is similar in many respects to zinc but forms more complex compounds.
Chemical The most common oxidation state of cadmium is +2, though rare examples of +1 can be found. Cadmium burns in air to form brown amorphous cadmium oxide (CdO). The crystalline form of the same compound is dark red and changes colour when heated, similar to zinc oxide. Hydrochloric acid, sulfuric acid and nitric acid dissolve cadmium by forming cadmium chloride (CdCl2) cadmium sulfate (CdSO4) or cadmium nitrate (Cd(NO3)2). The oxidation state +1 can be reached by dissolving cadmium in a mixture of cadmium chloride and aluminium chloride, forming the Cd22+ which is similar to the Hg22+ in mercury(I) chloride.[2] Cd + CdCl2 + 2 AlCl3 → Cd2[AlCl4]2
Isotopes Naturally occurring cadmium is composed of 8 isotopes. For two of them, natural radioactivity was observed, and three others are predicted to be radioactive but their decay is not observed, due to extremely long half-life times. The two natural radioactive isotopes are 113Cd (beta decay, half-life is 7.7 × 1015 years) and 116Cd (two-neutrino double beta decay, half-life is 2.9 × 1019 years). The other three are 106Cd, 108Cd (double electron capture), and 114Cd (double beta decay); only lower limits on their half-life times have been set. At least three isotopes - 110Cd, 111Cd, and 112Cd - are stable. Among the isotopes absent in natural cadmium, the most long-lived are 109Cd with a half-life of 462.6 days, and 115 Cd with a half-life of 53.46 hours. All of the remaining radioactive isotopes have half-lives that are less than 2.5 hours, and the majority of these have half-lives that are less than 5 minutes. This element also has 8 known meta states, with the most stable being 113m Cd (t½ 14.1 years), 115mCd (t½ 44.6 days), and 117mCd (t½ 3.36 hours). The known isotopes of cadmium range in atomic mass from 94.950 u (95Cd) to 131.946 u (132Cd). The primary decay mode before the second-most-abundant stable isotope, 112Cd, is electron capture, and the primary modes after are beta emission. The primary decay product before 112Cd is element 47 (silver), and the primary product after is element 49 (indium). One isotope of cadmium, 113Cd, absorbs neutrons with very high probability if they have an energy below the The cadmium-113 total cross section cadmium cut-off and transmits them readily otherwise. clearly showing the cadmium cutoff. The cadmium cut-off is about 0.5 eV.[3] Neutrons with energy below the cutoff are deemed slow neutrons, distinguishing them from intermediate and fast neutrons.
Cadmium
5
Applications Batteries About three-quarters of all the cadmium used is in batteries, predominantly in rechargeable nickel-cadmium batteries. Nickel-cadmium cells have a nominal cell potential of 1.2 V. The cell consists of a positive nickel hydroxide electrode and a negative cadmium electrode plate separated by an alkaline electrolyte (potassium hydroxide). More recent nickel-metal hydride batteries reduce the use of Ni-Cd Ni-Cd batteries batteries. The European Union banned the use of cadmium in electronics in 2004 with several exceptions but reduced the allowed content of cadmium in electronics to 0.002 %.[4]
Other uses Most of the remaining quarter is used mainly for cadmium pigments, coatings and plating, and as stabilizers for plastics. Other uses include: • In some of the lowest-melting alloys, for example Wood's metal • In bearing alloys, due to a low coefficient of friction and very good fatigue resistance[5] • In electroplating (6% cadmium)[5] . Cadmium electroplating is widely used in aircraft industry due to the excellent corrosion resistance of cadmium-plated steel components. Cadmium provides cathodic protection to low-alloyed steels, Train painted with cadmium yellow since it is positioned lower in the galvanic series. The coating is usually passivated by chromate salts. A significant limitation of cadmium plating is hydrogen embrittlement of high-strength steels caused by the electroplating process. Therefore, steel parts heat-treated to tensile strength above 1300 MPa (200 ksi) should be coated by an alternative method (such as special low-embrittlement cadmium electroplating processes or physical vapor deposition). • In many kinds of solder[5] • As a barrier to control neutrons in nuclear fission[5] • The pressurized water reactor designed by Westinghouse Electric Company uses an alloy consisting of 80% silver, 15% indium, and 5% cadmium.[5] • Cadmium oxide in black and white television phosphors and in the blue and green phosphors for colour television picture tubes[6] • Cadmium sulfide (CdS) as a photoconductive surface coating for photocopier drums. • In paint pigments, cadmium forms various salts, with CdS being the most common. This sulfide is used as a yellow pigment. Cadmium selenide can be used as red pigment, commonly called cadmium red. To painters who work with the pigment, cadmium yellows, oranges, and reds are the most potent colours to use. In fact, during production, these colours are significantly toned down before they are ground with oils and binders,
Cadmium
6
or blended into watercolours, gouaches, acrylics, and other paint and pigment formulations. These pigments are toxic, and it is recommended to use a barrier cream on the hands to prevent absorption through the skin when working with them. • In some semiconductors such as cadmium sulfide, cadmium selenide, and cadmium telluride, which can be used for light detection or solar cells. HgCdTe is sensitive to infrared. • In PVC as stabilizers. • In molecular biology, it is used to block voltage-dependent calcium channels from fluxing calcium ions. See also Category:Cadmium compounds.
History Cadmium (Latin cadmia, Greek καδμεία meaning "calamine", a cadmium-bearing mixture of minerals, which was named after the Greek mythological character, Κάδμος Cadmus) was discovered in Germany in 1817 by Friedrich Strohmeyer.[7] Strohmeyer found the new element as an impurity in zinc carbonate (calamine), and, for 100 years, Germany remained the only important producer of the metal. The metal was named after the Latin word for calamine, since the metal was found in this zinc compound. Strohmeyer noted that some impure samples of calamine changed colour when heated but pure calamine did not. He was persistent in studying these results and eventually isolated cadmium metal by roasting and reduction of the sulfide. Even though cadmium and its compounds are highly toxic, the British Pharmaceutical Codex from 1907 states that cadmium iodide was used as a medicine to treat "enlarged joints, scrofulous glands,[8] and chilblains".
Friedrich Strohmeyer
Cadmium
7 In 1927, the International Conference on Weights and Measures redefined the metre in terms of a red cadmium spectral line (1m = 1,553,164.13 [9] wavelengths). This definition has since been changed (see krypton).
After the industrial scale production of cadmium started in the 1930s and 1940s the major application was the coating of steel and copper alloys to prevent World production trend corrosion. In 1944 62 % and in 1956 59 % of the cadmium in the United States was used for this [10] purpose. The second application where red and yellow pigments based on sulfides and selenides of cadmium. In 1956 24 % of the cadmium used within the United States was used for this purpose.[10] The stabilizing effect of cadmium containing chemicals on plastics led to a increased use of those compounds in the 1970s and 1980s. The use of Cadmium in all applications mentioned above declined drastically due to environmental and health regulations from 1980 on. In 2006 only 7 % of the cadmium is used for plating and coating and only 10% is used for pigments. The decrease in the consumption in other applications was made up by a growing demand of cadmium in nickel-cadmium batteries, which accounted for 81% of the cadmium consumption in the United States in 2006. The overall consumption of cadmium was nearly constant from the 1970s till 2009.[10]
Occurrence Cadmium-containing ores are rare and are found to occur in small quantities. However, traces do naturally occur in phosphate, and have been shown to transmit in food through fertilizer application.[11] Greenockite (CdS), the only cadmium mineral of importance, is nearly always associated with sphalerite (ZnS). As a consequence, cadmium is produced mainly as a byproduct from mining, smelting, and refining sulfide ores of zinc, and, to a lesser degree, lead and copper. Cadmium metal Small amounts of cadmium, about 10% of consumption, are produced from secondary sources, mainly from dust generated by recycling iron and steel scrap. Production in the United States began in 1907, but it was not until after World War I that cadmium came into wide use. One place where metallic cadmium can be found is the Vilyuy River basin in Siberia.[12] See also Category:Cadmium minerals.
Cadmium
8
Extraction In 2001, China was the top producer of cadmium with almost one-sixth world share closely followed by South Korea and Japan, reports the British Geological Survey. Cadmium is a common impurity in zinc ores, and it is most often isolated during the production of zinc. Some zinc ores concentrates from sulfidic zinc ores contain Cadmium output in 2005 [13] up to 1,4 % of cadmium. In 1970s the output of cadmium was 6.5 pounds per ton of zinc.[13] Zinc sulfide ores are roasted in the presence of oxygen, converting the zinc sulfide to the oxide. Zinc metal is produced either by smelting the oxide with carbon or by electrolysis in sulfuric acid. Cadmium is isolated from the zinc metal by vacuum distillation if the zinc is smelted, or cadmium sulfate is precipitated out of the electrolysis solution.[14]
Biological role A role of cadmium in biology has been recently discovered. A cadmium-dependent carbonic anhydrase has been found in marine diatoms. Cadmium does the same job as zinc in other anhydrases, but the diatoms live in environments with very low zinc concentrations, thus biology has taken cadmium rather than zinc, and made it work. The discovery was made using X-ray absorption fluorescence spectroscopy (XAFS), and cadmium was characterized by noting the energy of the X-rays that were absorbed.[15] [16]
Toxicity Cadmium
poisoning
is
an
occupational
hazard
associated with industrial processes such as metal plating and the production of nickel-cadmium batteries, pigments, plastics, and other synthetics. The primary route of exposure in industrial settings is inhalation. Inhalation of cadmium-containing fumes can result initially in metal fume fever but may progress to chemical pneumonitis, pulmonary edema, and death.[17]
Image of the violet light from a helium
cadmium metal vapor laser. The highly Cadmium is also a potential environmental hazard. monochromatic color arises from the Human exposures to environmental cadmium are 441.563 nm transition line of cadmium. primarily the result of the burning of fossil fuels and municipal wastes.[18] However, there have been notable instances of toxicity as the result of long-term exposure to cadmium in contaminated food and water. In the decades leading up to World War II, Japanese mining operations contaminated the Jinzu River with cadmium and traces of other toxic metals. As a consequence, cadmium accumulated in the rice crops growing along the riverbanks downstream of the mines. The local agricultural communities consuming the contaminated rice developed Itai-itai disease and renal abnormalities, including proteinuria and glucosuria.[19] Cadmium is one of six substances banned by the European Union's Restriction on Hazardous Substances (RoHS) directive, which bans certain hazardous substances in electronics.
Cadmium Cadmium and several cadmium-containing compounds are known carcinogens and can induce many types of cancer.[20] Research has found that cadmium toxicity may be carried into the body by zinc binding proteins; in particular, proteins that contain zinc finger protein structures. Zinc and cadmium are in the same group on the periodic table, contain the same common oxidation state (+2), and when ionized are almost the same size. Due to these similarities, cadmium can replace zinc in many biological systems, in particular, systems that contain softer ligands such as sulfur. Cadmium can bind up to ten times more strongly than zinc in certain biological systems, and is notoriously difficult to remove. In addition, cadmium can replace magnesium and calcium in certain biological systems, although these replacements are rare. Tobacco smoking is the most important single source of cadmium exposure in the general population. It has been estimated that about 10% of the cadmium content of a cigarette is inhaled through smoking. The absorption of cadmium from the lungs is much more effective than that from the gut, and as much as 50% of the cadmium inhaled via cigarette smoke may be absorbed.[21] On average, smokers have 4-5 times higher blood cadmium concentrations and 2-3 times higher kidney cadmium concentrations than non-smokers. Despite the high cadmium content in cigarette smoke, there seems to be little exposure to cadmium from passive smoking. No significant effect on blood cadmium concentrations could be detected in children exposed to environmental tobacco smoke.[22]
See also • List of breast carcinogenic substances
External links • ATSDR Case Studies in Environmental Medicine: Cadmium Toxicity [23] U.S. Department of Health and Human Services • IARC Monograph "Cadmium and Cadmium Compounds" [24] • National Pollutant Inventory - Cadmium and compounds [25] • WebElements.com – Cadmium [26] • Los Alamos National Laboratory – Cadmium [27] • [28] Warning Moose and Deer Liver • National Institute for Occupational Safety and Health - Cadmium Page [29] • USGS Comodity Report cadmium [30]
9
Cadmium
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils; (1985) (in German). Lehrbuch der Anorganischen Chemie (91–100 ed.). Walter de Gruyter. pp. 1056–1057. ISBN 3-11-007511-3. [3] Knoll, G.F. (1999). Radiation Detection and Measurement, 3rd edition. Wiley. ISBN 978-0471073383. p505 [4] " Battery collection;recycling, nature protected (http:/ / www. europarl. europa. eu/ sides/ getDoc. do?pubRef=-/ / EP/ / TEXT+ IM-PRESS+ 20060628BRI09328+ FULL-TEXT+ DOC+ XML+ V0/ / EN)". . Retrieved 2008-11-04. [5] Scoullos, Michael J.; Vonkeman, Gerrit H.; Thornton, Iain; Makuch, Zen (2001). Mercury, Cadmium, Lead: Handbook for Sustainable Heavy Metals Policy and Regulation (http:/ / books. google. de/ books?id=9yzN-QGag_8C). Springer. ISBN 9781402002243. . [6] Lee, Ching-Hwa; Hs, Chi-Shiung (2002). "Recycling of Scrap Cathode Ray Tubes". Environmental Science and Technology 36 (1): 69–75. doi: 10.1021/es010517q S0013-936X(01)00517-X (http:/ / dx. doi. org/ 10. 1021/ es010517q+ S0013-936X(01)00517-X). [7] Hermann; Strohmeyer F. (1818). "Noch ein schreiben über das neue Metall(Another letter about the new metal)". Annalen der Physik 59: 113. [8] Dunglison, Robley (1866). Medical Lexicon: A Dictionary of Medical Science (http:/ / books. google. com/ books?id=PmohO5jV2YsC). Henry C. Lea. pp. 159. . [9] Burdun, G. D. (1958). " On the new determination of the meter (http:/ / www. springerlink. com/ content/ tk70442064438147/ fulltext. pdf?page=1)" (pdf). Measurement Techniques 1 (3): 259–264. doi: 10.1007/BF00974680 (http:/ / dx. doi. org/ 10. 1007/ BF00974680). . [10] Lansche, Arnold M.. " Minerals Yearbook 1956: Cadmium (http:/ / digicoll. library. wisc. edu/ cgi-bin/ EcoNatRes/ EcoNatRes-idx?type=turn& entity=EcoNatRes. MinYB1956v1. p0289& id=EcoNatRes. MinYB1956v1& isize=XL& q1=cadmium)". United States Geological Survey. . Retrieved 2008-04-21. [11] Jiao, You; Grant, Cynthia A.; Bailey, Loraine D.. "Effects of phosphorus and zinc fertilizer on cadmium uptake and distribution in flax and durum wheat". Journal of the Science of Food and Agriculture 84 (8): 777–785. doi: 10.1002/jsfa.1648 (http:/ / dx. doi. org/ 10. 1002/ jsfa. 1648). [12] Fleischer, Michael; Cabri, LouisJ.; Chao, Georg Y.; Papst, Adolf (1980). " New Mineral Names (http:/ / www. minsocam. org/ ammin/ AM65/ AM65_1065. pdf)". American Mineralogist 65: 1065–1070. . [13] National Research Council (U.S.), Panel on Cadmium, Committee on Technical Aspects of Critical and Strategic Material (1969). Trends in Usage of Cadmium: Report (http:/ / books. google. de/ books?id=okArAAAAYAAJ). National Research Council, National Academy of Sciences-National Academy of Engineering. pp. 1–3. . [14] Cadmium (http:/ / www. webelements. com/ webelements/ elements/ text/ Cd/ key. html) at WebElements.com [15] Lane, Todd W.; Morel, François M. M. (2000). " A biological function for cadmium in marine diatoms (http:/ / www. pnas. org/ content/ 97/ 9/ 4627. full. pdf+ html)". Proc. Natl. Acad. Sci. 97 (9): 4627–4631. doi: 10.1073/pnas.090091397 (http:/ / dx. doi. org/ 10. 1073/ pnas. 090091397). . [16] Lane, Todd W. ,; Saito, Mak A.; George, Graham N.; Pickering, Ingrid J.; Prince, Roger C.; Morel, François M. M. (2005). " A cadmium enzyme from a marine diatom (http:/ / www. whoi. edu/ cms/ files/ msaito/ 2005/ 5/ LaneSaitoMorel_CdCA_Nature2005_2944. pdf)". Nature 435 (42). doi: 10.1038/435042a (http:/ / dx. doi. org/ 10. 1038/ 435042a). . [17] Hayes, Andrew Wallace (2007). Principles and Methods of Toxicology (http:/ / books. google. de/ books?id=vgHXTId8rnYC& dq). Philadelphia: CRC Press. . [18] " EPA summary on cadmium (http:/ / www. epa. gov/ ttn/ atw/ hlthef/ cadmium. html)". U.S. Environmental Protection Agency. . Retrieved 2008-04-21. [19] Nogawa, Koji; Kobayashi, Etsuko; Okubo, Yasushiand; Suwazono, Yasushi (2004). " Environmental cadmium exposure, adverse effects, and preventative measures in Japan (http:/ / www. springerlink. com/ content/ n0773057mw738u05/ )". Biometals 17 (5): 581–587. doi: 10.1023/B:BIOM.0000045742.81440.9c (http:/ / dx. doi. org/ 10. 1023/ B:BIOM. 0000045742. 81440. 9c). . [20] " 11th Report on Carcinogens (http:/ / ntp. niehs. nih. gov/ index. cfm?objectid=32BA9724-F1F6-975E-7FCE50709CB4C932)". National Toxicology Program. . Retrieved 2008-04-21. [21] Friberg, L. (1983). "Cadmium". Annual Review of Public Health 4: 367–367. doi: 10.1146/annurev.pu.04.050183.002055 (http:/ / dx. doi. org/ 10. 1146/ annurev. pu. 04. 050183. 002055). [22] Jarup, L.; Berglund, M.; Elinder, C. G.; Nordberg, G.; Vahter, M (1998). "Health effects of cadmium exposure—a review of the literature and a risk estimate.". Scandinavian Journal of Work, Environment and Health 24: 11–51.
10
Cadmium [23] [24] [25] [26] [27] [28] [29] [30]
http:/ / www. atsdr. cdc. gov/ csem/ cadmium/ http:/ / www-cie. iarc. fr/ htdocs/ monographs/ vol58/ mono58-2. htm http:/ / www. npi. gov. au/ database/ substance-info/ profiles/ 17. html http:/ / www. webelements. com/ webelements/ elements/ text/ Cd/ index. html http:/ / periodic. lanl. gov/ elements/ 48. html http:/ / publicdocs. mnr. gov. on. ca/ View. asp?Document_ID=10690& Attachment_ID=25585 http:/ / www. cdc. gov/ niosh/ topics/ Cadmium http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ cadmium/
11
Article Sources and Contributors
Article Sources and Contributors Cadmium Source: http://en.wikipedia.org/w/index.php?oldid=308492267 Contributors: A Raider Like Indiana, Aadal, Aaronmz, Acadapter44, Adamrush, Ahoerstemeier, Ahruman, Alansohn, Alarichus, Anclation, AndonicO, Andros 1337, Animum, Anlace, Antandrus, Anwar saadat, Archimerged, Arkuat, AssegaiAli, Aussie Alchemist, Axiosaurus, Beano, Beetstra, Betacommand, Black and White, BlueEarth, Bobo192, Borislav.dopudja, Bork, Brado1001, Brian Huffman, BrianGV, Bryan Derksen, Burntsauce, CYD, Caltas, Can't sleep, clown will eat me, CanisRufus, Carnildo, Cbustapeck, Chester737, Chris 73, Clarenceville Trojan, Conversion script, Corvus cornix, Cryptic C62, Damnreds, Darrien, David Latapie, Ddroar, Deglr6328, Delta G, Deor, DesertAngel, Dlae, Donarreiskoffer, Drphilharmonic, Dufo, Dwmyers, Echris1, Edgar181, El C, Eldin raigmore, Emperorbma, Emurph, Epbr123, Esprit15d, EvilPettingZoo, Evolution686us, Excirial, Faradayplank, Feezo, Femto, Freeeedom, Fresheneesz, Frosted14, Gaius Cornelius, Geoff, Georgewilliamherbert, Gerry Ashton, Giftlite, Gimmetrow, GregorB, Grendelkhan, Hak-kâ-ngìn, Halfdan, Hede2000, Heron, Holtth, Icairns, Illustria, J.delanoy, JDspeeder1, JWSchmidt, JaGa, Jaan513, JadziaLover, JakeLucas, Jaraalbe, Jaxl, Jay Litman, Jeff-swanson, Jerzy, Joanjoc, John, JohnOwens, Johnleemk, Jose77, Jrockley, Julesd, Jumbuck, Kanjilearner, Karl-Henner, KarlJorgensen, Keith Edkins, Ken Arromdee, Khoptiar il, Koyaanis Qatsi, Kukini, Kungming2, Kurykh, Kwamikagami, Leebo, Lethalgeek, Lewis R, Leyo, Linuxmatt, Lithpiperpilot, Lordryker, MONGO, Mandel, Mashford, Mauimonica, Mav, Mendaliv, Michael Hardy, Minesweeper, Mr0t1633, NE2, Nakon, Nergaal, Nihiltres, Ntouran, NucleaMachina, Od Mishehu, OwenX, Oxymoron83, Paul August, Peridon, Peyre, Physchim62, Piperh, Plantsurfer, Polly, Poolkris, Pras, Proguitar, Pusher, Quadrius, QueenCake, Quinsareth, RTC, Raul654, Redux, Remember, Rhelmich, Riana, Roberta F., Romanm, Ryulong, Samuell, Saperaud, Schneelocke, Seanisaweird, Sengkang, Shafei, Shanes, Skatebiker, Sl, Smallweed, Spens10, Spiritia, Squids and Chips, Stefan da, Stephenb, Stifynsemons, Still, Stone, Stud muffin09, Suisui, Sunborn, SuperDude115, Tagishsimon, Taranah, Tetracube, Thecurran, Theseeker4, Thingg, Thricecube, Timeastor, Tisdalepardi, Titoxd, Tomaxer, Tristanb, Tullywinters, Uusitunnus, V1adis1av, VermillionBird, Versus22, Vsmith, Vuong Ngan Ha, Waggers, Waisberg, Warut, Watch37264, Waterspell1, Wavelength, WhatamIdoing, Wikiscient, Yalens, Yekrats, Yyy, 308 anonymous edits
Image Sources, Licenses and Contributors image:Cd-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Cd-TableImage.png License: GNU Free Documentation License Contributors: Conscious, Mav, Paddy, Paginazero, Saperaud Image: Cd,48.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Cd,48.jpg License: GNU Free Documentation License Contributors: User:RTC Image:Cadmium cutoff.png Source: http://en.wikipedia.org/w/index.php?title=File:Cadmium_cutoff.png License: Public Domain Contributors: M.B. Chadwick, P. Oblozinsky, M. Herman at al. Image:NiCd various.jpg Source: http://en.wikipedia.org/w/index.php?title=File:NiCd_various.jpg License: GNU Free Documentation License Contributors: ARTE, Boffy b, Glenn, HenkvD Image:Tyne and Wear Metro train 4001 at Pelaw 01.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Tyne_and_Wear_Metro_train_4001_at_Pelaw_01.jpg License: unknown Contributors: Jed, Thryduulf File:Friedrich Strohmeyer.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Friedrich_Strohmeyer.jpg License: Public Domain Contributors: User Grzes14 on pl.wikipedia Image:Cadmium - world production trend.svg Source: http://en.wikipedia.org/w/index.php?title=File:Cadmium_-_world_production_trend.svg License: Public Domain Contributors: User:Leyo Image:CadmiumMetalUSGOV.jpg Source: http://en.wikipedia.org/w/index.php?title=File:CadmiumMetalUSGOV.jpg License: Public Domain Contributors: Original uploader was User: at en.wikipedia Image:2005cadmium.PNG Source: http://en.wikipedia.org/w/index.php?title=File:2005cadmium.PNG License: Public Domain Contributors: User:Anwar_saadat/bubble_maps_(FAQ) Image:HeCd laser.jpg Source: http://en.wikipedia.org/w/index.php?title=File:HeCd_laser.jpg License: GNU Free Documentation License Contributors: Deglr6328, 1 anonymous edits
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
12
Indium
1
Indium 49
cadmium ← indium → tin
Ga ↑
In ↓
Tl Periodic Table - Extended Periodic Table
General Name, Symbol, Number
indium, In, 49
Element category
poor metals
Group, Period, Block
13, 5, p
Appearance
silvery lustrous gray
Standard atomic weight
114.818(3) g·mol
Electron configuration
[Kr] 4d
−1
10
Electrons per shell
2
1
5s 5p
2, 8, 18, 18, 3 Physical properties
Phase
solid
Density (near r.t.)
7.31 g·cm−3
Liquid density at m.p.
7.02 g·cm−3
Melting point
429.7485 K (156.5985 °C, 313.8773 °F)
Boiling point
2345 K (2072 °C, 3762 °F)
Heat of fusion
3.281 kJ·mol−1
Heat of vaporization
231.8 kJ·mol−1
Specific heat capacity
(25 °C) 26.74 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1196
1325
1485
1690
1962
2340
Atomic properties
Indium
2
Crystal structure
tetragonal
Oxidation states
3, 2, 1 (amphoteric oxide)
Electronegativity
1.78 (Pauling scale)
Ionization energies (more)
1st: 558.3 kJ·mol−1 2nd: 1820.7 kJ·mol−1 3rd: 2704 kJ·mol−1
Atomic radius
167 pm
Covalent radius
142±5 pm
Van der Waals radius
193 pm Miscellaneous [1]
Magnetic ordering
diamagnetic
Electrical resistivity
(20 °C) 83.7 n Ω·m
Thermal conductivity
(300 K) 81.8 W·m
Thermal expansion
(25 °C) 32.1 µm·m
Speed of sound (thin rod)
(20 °C) 1215 m/s
Young's modulus
11 GPa
−1
−1
·K
−1
Mohs hardness
−1
·K
1.2
Brinell hardness
8.83 MPa
CAS registry number
7440-74-6 Most-stable isotopes
Main article: Isotopes of indium iso
NA
half-life
113
4.3%
113
115
95.7%
4.41×1014 y
In In
DM
DE (MeV)
DP
In is stable with 64 neutron β−
0.495
115
Sn
References
Indium (pronounced /ˈɪndiəm/) is a chemical element with chemical symbol In and atomic number 49. This rare, soft, malleable and easily fusible post-transition metal is chemically similar to aluminium or gallium but more closely resembles zinc (zinc ores are also the primary source of this metal). Indium's current primary application is to form transparent electrodes from indium tin oxide in liquid crystal displays and touchscreens, and this use largely determines its global mining production. It is widely used in thin-films to form lubricated layers (during World War II it was widely used to coat bearings in high-performance aircraft). It is also used for making particularly low melting point alloys, and is a component in some lead-free solders. Radioactive indium-111 is used in nuclear medicine as a imaging agent to follow the movement of leukocytes in the body.
Indium
3
Characteristics Indium is a very soft, silvery-white, relatively rare true metal with a bright luster. As a pure metal indium emits a high-pitched "cry", when it is bent.[2] Both gallium and indium are able to wet glass.
Indium wetting the glass surface of a test tube
One unusual property of indium is that its most common isotope is slightly radioactive; it very slowly decays by beta emission to tin. This radioactivity has a half-life of 4.41 × 1014 years, four orders of magnitude larger than the age of the universe and nearly 50,000 times longer than that of natural thorium. Unlike its period 5 neighbor cadmium, indium is not a cumulative
poison.
Applications The first large-scale application for indium was as a coating for bearings in high-performance aircraft engines during World War II. Afterward, production gradually increased as new uses were found in fusible alloys, solders, and electronics. In the 1950s, tiny beads of it were used for the emitters and collectors of PNP alloy junction transistors. In the middle and late 1980s, the development of indium phosphide semiconductors and indium tin oxide thin films for liquid crystal displays (LCD) aroused much interest. By 1992, the thin-film application had become the largest end use.[3] [4]
Electronics • Indium oxide (In2O3) and indium tin oxide (ITO) are used as a transparent conductive coating applied to glass substrates in the making of electroluminescent panels. • Some indium compounds such as indium antimonide, indium phosphide,[5] and indium nitride[6] are semiconductors with useful properties. • Indium is used in the synthesis of the semiconductor copper indium gallium selenide (CIGS), which is used for the manufacture of thin film solar cells.[7] • Used in light-emitting diodes (LEDs) and Laser Diodes (LDs) based on compound semiconductors such as InGaN, InGaP that are fabricated by Metalorganic Vapor Phase Epitaxy (MOVPE) technology. • The ultrapure metalorganics of indium, specifically high purity trimethylindium (TMI) is used as a precursor in III-V compound semiconductors, while it is also used as the semiconductor dopant in II-VI compound semiconductors.[8]
Indium
Alloys and Metal • For manufacture of low-melting-temperature alloys. An alloy consisting of 24% indium and 76% gallium is liquid at room temperature. • Can also be plated onto metals and evaporated onto glass which forms a mirror which is as good as those made with silver[9] but has higher corrosion resistance. • Very small amounts used in aluminium alloy sacrificial anodes (for salt water applications) to prevent passivation of the aluminium. • To bond gold electrical test leads to superconductors, Indium is used as a conducting glue and applied under a microscope with precision tweezers. • In the form of a wire it is used as a vacuum seal in cryogenics and ultra-high vacuum applications. • Used as a calibration material for Differential scanning calorimetry.
Other uses • Indium tin oxide is used as a light filter in low pressure sodium vapor lamps. The infrared radiation is reflected back into the lamp, which increases the temperature within the tube and therefore improves the performance of the lamp.[4] • Indium's melting point of 429.7485 K (156.5985 °C) is a defining fixed point on the international temperature scale ITS-90. • Indium's high neutron capture cross section for thermal neutrons makes it suitable for use in control rods for nuclear reactors, typically in an alloy containing 80% silver, 15% indium, and 5% cadmium. • In nuclear engineering, the (n,n') reactions of 113In and 115In are used to determine magnitudes of neutron fluxes. • 111In emits gamma radiation and is used in scintigraphy, a technique of medical imaging. Scintigraphy has many applications, including early phase drug development, and monitoring the activity of white blood cells. A blood test is taken from the patient, white cells removed and labeled with the radioactive 111In, then re-injected back into the patient. Gamma imaging will reveal any areas of high white cell activity such as an abscess. • Indium is also used as a thermal interface material by personal computer enthusiasts in the form of pre-shaped foil sheets fitted between the heat-transfer surface of a microprocessor and its heat sink. The application of heat partially melts the foil and allows the indium metal to fill in any microscopic gaps and pits between the two surfaces, removing any insulating air pockets that would otherwise compromise heat transfer efficiency.
History The German chemists Ferdinand Reich and Hieronymous Theodor Richter were testing ores from the mines around Freiberg, Saxony. They dissolved the minerals pyrite, arsenopyrite, galena and sphalerite in hydrochloric acid and distilled the raw zinc chloride. As it was known that ores from that region sometimes contain thallium they searched for the green emission lines with spectroscopic methods. The green lines were absent but a blue line was present in the spectrum. As no element was known with a bright blue emission they concluded that a new element was present in the minerals. They named the element with the blue spectral line indium.[10] [11] Richter went on to isolate the metal in 1864.[12]
4
Indium
5
Until 1924, only approximately a gram of indium constituted the world's supply.
Occurrence and consumption Indium ranks 61st in abundance in the Earth's crust at approximately 0.25 ppm,[13] which means it is more than three times as abundant as silver, which occurs at 0.075 ppm.[14] Fewer than 10 indium minerals are known, none occurring in significant deposits. Examples are the dzhalindite (In(OH)3) and indict (FeIn2S4).[15]
Ductile Indium wire
Production The lack of indium mineral deposits and the fact that
indium is enriched in sulfidic lead, tin, copper, iron and predominately in zinc deposits, makes the zinc production the main source for indium. The indium is leached from slag and dust of zinc production. Further purification is done by electrolysis.[16] Up until 1924, there was only about a gram of isolated indium on the planet. Indium is produced mainly from residues generated during zinc ore processing but is also found in iron, lead, and copper ores.[2] Canada is a leading producer of indium. The Teck Cominco refinery in Trail, British Columbia, is the largest single source, with production of 32,500 kg in 2005, 41,800 kg in 2004 and 36,100 kg in 2003. The amount of indium consumed is largely a function of worldwide LCD production. Worldwide production is currently 476 tonnes per year from mining and a further 650 tonnes per year from recycling.[17] Demand has risen rapidly in recent years with the popularity of LCD computer monitors and televisions, which now account for 50% of indium consumption.[18] Increased manufacturing efficiency and recycling (especially in Japan) maintain a balance between demand and supply. Demand increased as the metal is used in LCDs and televisions, and supply decreased when a number of Chinese mining concerns stopped extracting indium from their zinc tailings. In 2002, the price was US$94 per kilogram. The recent changes in demand and supply have resulted in high and fluctuating prices of indium, which from 2005 to 2007 ranged from US$700/kg to US$1,000/kg.[19] Demand for indium may TFT display increase with large-scale manufacture of CIGS-based thin film solar technology starting by several companies in 2008, including Nanosolar and Miasole, although zinc oxide is often used instead.
Indium
6
Resources Based on content of indium in zinc ore stocks, there is a worldwide reserve base of approximately 6,000 tonnes of economically-viable indium.[19] This figure has led to estimates suggesting that, at current consumption rates, there is only 13 years' supply of indium left.[20] However, the Indium Corporation, the largest processor of indium, claims that, on the basis of increasing recovery yields during extraction, recovery from a wider range of base metals (including tin, copper and other polymetallic deposits) and new mining investments, the long-term supply of indium is sustainable, reliable and sufficient to meet increasing future demands.[17] This conclusion also seems reasonable in light of the fact that silver, a less abundant element, is currently mined at approximately 18,300 tonnes per annum,[21] which is 40 times greater than current indium mining rates. On the other hand, replacements for indium tin oxide are already on the horizon. According to recent research, mass production of transparent conductors made from graphene, a modification of the virtually inexhaustible element carbon discovered in 2004, may be just years away.[22]
Precautions Pure indium in metal form is considered non-toxic by most sources. In the welding and semiconductor industries, where indium exposure is relatively high, there have been no reports of any toxic side-effects. This may not be the case with indium compounds. For example, an anhydrous indium trichloride (InCl3) is quite toxic, while indium phosphide (InP) is both toxic and a suspected carcinogen.[23] [24]
See also • Indium compounds
References Notations • Los Alamos National Laboratory – Indium
[25]
External links • • • •
WebElements.com – Indium [26] The Indium Corporation [27] Reducing Agents > Indium low valent [28] Indium picture in the element collection from Heinrich Pniok
pnb:میڈنا
[29]
Indium
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] Alfantazi, A. M.; Moskalyk, R. R. (2003). "Processing of indium: a review". Minerals Engineering 16 (8): 687–694. doi: 10.1016/S0892-6875(03)00168-7 (http:/ / dx. doi. org/ 10. 1016/ S0892-6875(03)00168-7). [3] Tolcin, Amy C.. " Mineral Yearbook 2007: Indium (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ indium/ indiumyb06. pdf)" (pdf). United States Geological Survey. . Retrieved 2008-10-28. [4] Downs, Anthony John (1993). " Chemistry of Aluminium, Gallium, Indium, and Thallium (http:/ / books. google. com/ books?id=v-04Kn758yIC)". Springer. pp. 89 and 106. . [5] Bachmann, K. J. (1981). "Properties, Preparation, and Device Applications of Indium Phosphide". Annual Review of Materials Science 11: 441–484. doi: 10.1146/annurev.ms.11.080181.002301 (http:/ / dx. doi. org/ 10. 1146/ annurev. ms. 11. 080181. 002301). [6] Bhuiyan, Ghani; Hashimoto, Akihiro; Yamamoto, Akioare (2003). "Indium nitride (InN): A review on growth, characterization, and properties". Journal of Applied Physics 94: 2779. doi: 10.1063/1.1595135 (http:/ / dx. doi. org/ 10. 1063/ 1. 1595135). [7] Powalla, M.; Dimmler, B. (2000). "Scaling up issues of CIGS solar cells". Thin Solid Films 361-362: 540–546. doi: 10.1016/S0040-6090(99)00849-4 (http:/ / dx. doi. org/ 10. 1016/ S0040-6090(99)00849-4). [8] Shenai, Deodatta V.; Timmons, Michael L.; DiCarlo Jr., Ronald L.; Marsman, Charles J. (2004). "Correlation of film properties and reduced impurity concentrations in sources for III/V-MOVPE using high-purity trimethylindium and tertiarybutylphosphine". Journal of Crystal Growth 272: 603–608. doi: 10.1016/j.jcrysgro.2004.09.006 (http:/ / dx. doi. org/ 10. 1016/ j. jcrysgro. 2004. 09. 006). [9] Borra, E. F.; Tremblay, G.; Huot, Y.; Gauvin, J.. Publications of the Astronomical Society of the Pacific 109: 319-325. http:/ / adsabs. harvard. edu/ full/ 1997PASP. . 109. . 319B Title: Gallium Liquid Mirrors: Basic Technology, Optical-Shop Tests, and Observations. [10] Reich, F.; Richter, T. (1863). "Ueber das Indium". Journal für Praktische Chemie 90 (1): 172–176. doi: 10.1002/prac.18630900122 (http:/ / dx. doi. org/ 10. 1002/ prac. 18630900122). [11] Venetskii, S. (1971). "Indium". Metallurgist 15 (2): 148–150. doi: 10.1007/BF01088126 (http:/ / dx. doi. org/ 10. 1007/ BF01088126). [12] Reich, F.; Richter, T. (1864). "Ueber das Indium". Journal für Praktische Chemie 92 (1): 480–485. doi: 10.1002/prac.18640920180 (http:/ / dx. doi. org/ 10. 1002/ prac. 18640920180). [13] " The Element Indium (http:/ / education. jlab. org/ itselemental/ ele049. html)". It's Elemental. . Retrieved 2007-12-26. [14] " The Element Silver (http:/ / education. jlab. org/ itselemental/ ele047. html)". It's Elemental. . Retrieved 2007-12-26. [15] Sutherland, J. K. (01 Jun 1971). " A second occurrence of dzhalindite (http:/ / canmin. geoscienceworld. org/ cgi/ content/ abstract/ 10/ 5/ 781)". The Canadian Mineralogist 10 (5): 781–786. . [16] Schwarz-Schampera, Ulrich; Herzig, Peter M. (2002). Indium: Geology, Mineralogy, and Economics (http:/ / books. google. de/ books?hl=de& lr=& id=k7x_2_KnupMC). Springer. ISBN 9783540431350. . [17] " Indium and Gallium Supply Sustainability September 2007 Update (http:/ / www. indium. com/ _dynamo/ download. php?docid=552)" (pdf). 22nd EU PV Conference, Milan, Italy. . Retrieved 2007-12-26. [18] " Indium Price Supported by LCD Demand and New Uses for the Metal (http:/ / geology. com/ articles/ indium. shtml)" (pdf). Geology.com. . Retrieved 2007-12-26. [19] " Mineral Commodities Summary 2007: Indium (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ indium/ indiumcs07. pdf)" (pdf). United States Geological Survey. . Retrieved 2007-12-26. [20] " How Long Will it Last? (http:/ / environment. newscientist. com/ channel/ earth/ mg19426051. 200-earths-natural-wealth-an-audit. html)". New Scientist 194 (2605): 38–39. May 26, 2007. ISSN 0262-4079 (http:/ / worldcat. org/ issn/ 0262-4079). . [21] " Top World Silver Producers (http:/ / www. nma. org/ pdf/ g_silver_producers. pdf)" (pdf). World Silver Survey 2007. . [22] " Graphene-based gadgets may be just years away (http:/ / www. physorg. com/ news128776023. html)". June 23, 2008. . [23] Tanaka, A.; Hirata, M.; Omura, M., (2002). "Pulmonary toxicity of indium-tin oxide and indium phosphide after intratracheal instillations into the lung of hamsters". Journal of the Occupational Health 44: 99–102. doi: 10.1539/joh.44.99 (http:/ / dx. doi. org/ 10. 1539/ joh. 44. 99). [24] Blazka, ME; Dixon, D., Haskins, E., Rosenthal, G. J. (1994). "Pulmonary toxicity to intratracheally administered indium trichloride in Fischer 344 rats". Fundamental Applied Toxicology 22: 231–239. doi: 10.1006/faat.1994.1027 (http:/ / dx. doi. org/ 10. 1006/ faat. 1994. 1027). [25] http:/ / periodic. lanl. gov/ elements/ 49. html
7
Indium [26] [27] [28] [29]
8 http:/ / www. webelements. com/ webelements/ elements/ text/ In/ index. html http:/ / www. indium. com http:/ / www. organic-chemistry. org/ chemicals/ reductions/ indiumlowvalent. shtm http:/ / www. pse-mendelejew. de/ bilder/ in. jpg
Article Sources and Contributors
Article Sources and Contributors Indium Source: http://en.wikipedia.org/w/index.php?oldid=308238682 Contributors: A2Kafir, ABF, Aadal, Access, Ahoerstemeier, Alansohn, Anclation, Antandrus, Archimerged, Arkuat, Army1987, Arthana, AussieBoy, Avidallred, AxelBoldt, Baccyak4H, Bassplr19, Beetstra, Benbest, BlueEarth, Bomac, Bovineone, Brandonsgaywiki, Brockert, Bryan Derksen, BryanC, CYD, Caltas, Carnildo, Chris 73, Closedmouth, Conversion script, Cquan, CyrilB, DAID, Daniel.barna, Darrien, David Latapie, Dbo789, Delta G, DocWatson42, Dschwen, Eaolson, Edgar181, El C, Emperorbma, Endymi0n, Enok Walker, Epbr123, Erik Zachte, Femto, Gaius Cornelius, Gasheadsteve, Gene Nygaard, Gobeirne, Grendelkhan, Hak-kâ-ngìn, Hannibal, Hede2000, Helge Skjeveland, Herbee, Heron, II MusLiM HyBRiD II, IW.HG, Icairns, Imroy, J.delanoy, Jaan513, Jaraalbe, Jawed, Jeronimo, Joanjoc, John, JohnCub, Jopusbob, Jose77, Julesd, Karl23, Karlhahn, Katalaveno, Kwamikagami, LA2, LallLallLall, LarryMorseDCOhio, Loren.wilton, Mancunion, Mav, Maximus Rex, Mdinger, Mdwh, Mgimpel, Michael Hardy, Michaelbusch, Midgley, Mikeblas, Minesweeper, Miremare, Mm40, Momme, Mortdefides, NaBUru38, Nakon, Nergaal, Nihiltres, No Guru, NuclearWarfare, Nutiketaiel, Nweiheng, Oxymoron83, Paisa, Picapica, PierceG, Piperh, Pixel ;-), Plantsurfer, PlatinumX, Plexust, Poolkris, Pras, Pyrochem, RTC, Reedy, Remember, Reyk, Reza kalani, Rich Farmbrough, Rickhdz, Rifleman 82, Roberta F., Romanm, Saperaud, Sbharris, Schneelocke, Seeaxid, Sengkang, Sfuerst, ShaunMacPherson, SheepNotGoats, Shell Kinney, Sillybilly, Sl, Sloppy, Smack, Squids and Chips, Stephenb, Stifynsemons, Stone, Suisui, Tagishsimon, Tetracube, Thricecube, Traal, Trabert, Twinscimitars, Vsmith, Warut, Wiki alf, Wolfkeeper, Worldburns, Yartamis, Yuckfoo, Yyy, 212 anonymous edits
Image Sources, Licenses and Contributors image:In-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:In-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Mav, Paddy, Saperaud Image: Indium.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Indium.jpg License: unknown Contributors: User:Nerdtalker Image:Indium wetting glass.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Indium_wetting_glass.jpg License: Public Domain Contributors: User:Schtone Image:Indium wire.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Indium_wire.jpg License: unknown Contributors: User:Dschwen Image:Dell axim LCD under microscope.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Dell_axim_LCD_under_microscope.jpg License: GNU Free Documentation License Contributors: 1 anonymous edits
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9
Tin
1
Tin 50
indium ← tin → antimony
Ge ↑
Sn ↓
Pb Periodic Table - Extended Periodic Table
General Name, Symbol, Number
tin, Sn, 50
Element category
poor metals
Group, Period, Block
14, 5, p
Appearance
silvery lustrous gray
Standard atomic weight
118.710(7) g·mol
Electron configuration
[Kr] 4d
Electrons per shell
−1
10
2
2
5s 5p
2, 8, 18, 18, 4 Physical properties
Phase
solid
Density (near r.t.)
(white) 7.365 g·cm−3
Density (near r.t.)
(gray) 5.769 g·cm−3
Liquid density at m.p.
6.99 g·cm−3
Melting point
505.08 K (231.93 °C, 449.47 °F)
Boiling point
2875 K (2602 °C, 4716 °F)
Heat of fusion
(white) 7.03 kJ·mol−1
Heat of vaporization
(white) 296.1 kJ·mol−1
Specific heat capacity
(25 °C) (white) 27.112 J·mol−1·K−1
Tin
2
Vapor pressure P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1497
1657
1855
2107
2438
2893
Atomic properties Crystal structure
Tetragonal (white), diamond cubic (gray)
Oxidation states
4, 2, -4 (amphoteric oxide)
Electronegativity Ionization energies (more)
1.96 (Pauling scale) 1st: 708.6 kJ·mol−1 2nd: 1411.8 kJ·mol−1 −1
3rd: 2943.0 kJ·mol Atomic radius
140 pm
Covalent radius
139±4 pm
Van der Waals radius
217 pm Miscellaneous
Magnetic ordering Electrical resistivity
[1]
(gray)diamagnetic
, (white) paramagnetic
(0 °C) 115 nΩ·m
Thermal conductivity
(300 K) 66.8 W·m
Thermal expansion
(25 °C) 22.0 µm·m
Speed of sound (thin rod)
(r.t.) (rolled) 2730 m·s
Young's modulus
50 GPa
Shear modulus
18 GPa
Bulk modulus
58 GPa
−1
−1
0.36
Mohs hardness
1.5
CAS registry number
−1
·K
−1
Poisson ratio
Brinell hardness
−1
·K
51 MPa 7440-31-5 Most-stable isotopes
Tin
3
Main article: Isotopes of tin iso
NA
half-life
112
0.97%
112
114
0.66%
114
115
0.34%
115
116
14.54%
116
117
7.68%
117
118
24.22%
118
119
8.59%
119
120
32.58%
120
122
4.63%
122
124
5.79%
124
Sn Sn Sn Sn Sn Sn Sn Sn Sn Sn
126
Sn
syn
DM
DE (MeV)
DP
Sn is stable with 62 neutron Sn is stable with 64 neutron Sn is stable with 65 neutron Sn is stable with 66 neutron Sn is stable with 67 neutron Sn is stable with 68 neutron Sn is stable with 69 neutron Sn is stable with 70 neutron Sn is stable with 72 neutron Sn is stable with 74 neutron 5
~1×10 y
−
β
0.380
126
Sb
References
Tin is a chemical element with the symbol Sn (Latin: Stannum) and atomic number 50. It is a main group metal in group 14 of the periodic table. Tin shows chemical similarity to both neighboring group 14 elements, germanium and lead, like the two possible oxidation states +2 and +4. Tin is the 49th most abundant element and has, with 10 isotopes, the largest number of stable isotopes in the periodic table. Tin is obtained chiefly from the mineral cassiterite, where it occurs as tin dioxide, SnO2. This silvery, malleable poor metal is not easily oxidized in air, and is used to coat other metals to prevent corrosion. The first alloy used in large scale since 3000 BC was bronze, an alloy of tin and copper. After 600 BC pure metallic tin was produced. Pewter, which is an alloy of 85 % to 90 % tin with the remainder commonly consisting of copper, antimony and lead, was used for flatware from the Bronze Age until the 20th century. In modern times tin is used in many alloys, most notably tin/lead soft solders, typically containing 60% or more of tin. Another large application for tin is corrosion-resistant tin plating of steel. Due to its low toxicity, tin-plated metal is also used for food packaging, giving the name to tin cans, which are made mostly out of aluminium or tin-plated steel.
Characteristics Physical and allotropes Tin is a malleable, ductile, and highly crystalline silvery-white metal. It is malleable at ordinary temperatures but is brittle when it is cooled, due to the properties of its two major allotropes, α- and β-tin. When a bar of tin is bent, a crackling sound known as the tin cry can be heard due to the twinning of the crystals.[2] The two allotropes that are encountered at normal pressure and temperature, α-tin and β-tin, are more commonly known as gray tin,
Tin
4 and respectively white tin. Two more allotropes, γ and σ, exist at temperatures above 161 °C and pressures above several GPa.[3] White tin, or the β-form, is metallic, and is the stable one at room conditions or at higher temperatures. Below 13.2 °C, tin exists in the gray α-form, which has a diamond cubic crystal structure, similar to diamond, silicon or germanium. Gray tin has no metallic properties at all, is a dull-gray powdery material, and has few uses, other than a few specialized semiconductor applications.[2] Although the α-β transformation temperature is nominally 13.2 °C, impurities (e.g. Al, Zn, etc.) lower the transition temperature well below 0 °C, and upon addition of Sb or Bi the transformation may not occur at all.[4] This conversion is known as tin disease or tin pest. Tin pest was a particular problem in northern Europe in the 18th century as organ pipes made of tin alloy would sometimes be affected during long cold winters. Some sources also say that during Napoleon's Russian campaign of 1812, the temperatures became so cold that the tin buttons on the soldiers' uniforms disintegrated, contributing to the defeat of the Grande Armée. The veracity of this story is debatable, because the transformation to gray tin often takes a reasonably long time.[5] Commercial grades of tin (99.8%) resist transformation because of the inhibiting effect of the small amounts of bismuth, antimony, lead, and silver present as impurities. Alloying elements such as copper, antimony, bismuth, cadmium, and silver increase its hardness. Tin tends rather easily to form hard, brittle intermetallic phases, which are often undesirable. It does not form wide solid solution ranges in other metals in general, and there are few elements that have appreciable solid solubility in tin. Simple eutectic systems, however, occur with bismuth, gallium, lead, thallium, and zinc.[4]
Chemistry and compounds See also Tin compounds Tin is classified as a semimetal, as its chemical properties fall between those of metals and non-metals, just as the semiconductors silicon and germanium do. It resists corrosion from distilled, sea and soft tap water, but can be attacked by strong acids, alkalis, and acid salts. Tin can be highly polished and is used as a protective coat for other metals in order to prevent corrosion or other chemical action. Tin acts as a catalyst when oxygen is in solution and helps accelerate chemical attack.[2] Tin forms the dioxide SnO2 (cassiterite) when it is heated in the presence of air. SnO2, in turn, is feebly acidic and forms stannate (SnO32−) salts with basic oxides. There are also stannates with the structure [Sn(OH)6]2−, like K2[Sn(OH)6], although the free stannic acid H2[Sn(OH)6] is unknown. Tin combines directly with chlorine forming tin(IV) chloride, while reacting tin with hydrochloric acid in water gives tin(II) chloride and hydrogen. Several other compounds of tin exist in the +2 and +4 oxidation states, such as tin(II) sulfide and tin(IV) sulfide (Mosaic gold). There is only one stable hydride, however: stannane (SnH4), where tin is in the +4 oxidation state.[2] The most important salt is stannous chloride, which has found use as a reducing agent and as a mordant in the calico printing process. Electrically conductive coatings are produced when tin salts are sprayed onto glass. These coatings have been used in panel lighting and in the production of frost-free windshields. Tin is added to some dental care products[6] [7] as stannous fluoride (SnF2). Stannous fluoride can be mixed with calcium abrasives while the more common sodium fluoride
Tin
5 gradually becomes biologically inactive combined with calcium.[8] It has also been shown to be more effective than sodium fluoride in controlling gingivitis.[9] Organotin compounds or stannanes are chemical compounds based on tin with hydrocarbon substituents.[10] Organotin compounds usually have high toxicity and have been used as biocides, but their use is slowly being phased out. The first organotin compound was diethyltin diiodide (Sn(C2H5)2I2), discovered by Edward Frankland in 1849. Organotin compounds differ from their lighter analogues of germanium and silicon in that there is a greater occurrence of the +2 oxidation state due to the "inert pair effect"; it also has a greater range of coordination numbers, and the common presence of halide bridges between polynuclear compounds. Most organotin compounds are colorless liquids or solids that are usually stable to air and water. The tetraalkyl stannates (R4Sn) always have a tetrahedral geometry at the tin atom. The halide derivatives R3SnX often form chained structures with Sn-X-Sn bridges. Alkyltin compounds are usually prepared via Grignard reagent reactions such as in: SnCl4 + 4 RMgBr → R4Sn + 4 MgBrCl.[11]
Isotopes Tin is the element with the greatest number of stable isotopes, ten; these include all those with atomic masses between 112 and 124, with the exception of 113, 121 and 123. Of these, the most abundant ones are 120Sn (at almost a third of all tin), 118Sn, and 116Sn, while the least abundant one is 115Sn. The isotopes possessing even atomic numbers have no nuclear spin while the odd ones have a spin of +1/2. Tin, with its three common isotopes 115Sn, 117 Sn and 119Sn, is among the easiest elements to detect and analyze by NMR spectroscopy, and its chemical shifts are referenced against SnMe4.[12] [13] This large number of stable isotopes is thought to be a direct result of tin possessing an atomic number of 50, which is a "magic number" in nuclear physics. There are 28 additional unstable isotopes that are known, encompassing all the remaining ones with atomic masses between 99 and 137. Aside from 126Sn, which has a half-life of 230,000 years, all the radioactive isotopes have a half-life of less than a year. The radioactive 100Sn is one of the few nuclides possessing a "doubly magic" nucleus and was discovered relatively recently, in 1994.[14] Another 30 metastable isomers have been characterized for isotopes between 111 and 131, the most stable of which being 121mSn, with a half-life of 43.9 years.
Etymology The Latin name Stannum is connected to "stagnum" and "stag" (Indo-European) for dripping because tin melts easily. The former "stagnum" was the word for a stale pool or puddle, with a cognate in the English word "stagnant." The English word "tin" has cognates in many Germanic and Celtic languages. The American Heritage Dictionary speculates that the word was borrowed from a pre-Indo-European language. The later name "stannum" and its Romance derivatives come from the lead-silver alloy of the same name for the finding of the silver in ores. The word definitely assumed its present meaning in the 4th century (H. Kopp). According to Meyers Konversationslexikon Stannum is derived from Cornish stean (present orthography sten), and is proof that Cornwall in the first centuries AD was the main source of tin. Other sources, however, see the Cornish stean merely as a back-derivation from the Latin stannum [15] . The Latin Stannum became the source for most European words.
Tin
6 According to SMI [16] the English word for the metal is named after an Etruscan god, Tinia. (variants include Old English: tin, Old Latin: plumbum candidum ("white lead"), Old German: tsin, Late Latin: stannum)
History Antiquity Tin is one of the earliest metals known.[17] Late Stone Age metal-workers discovered that putting a small amount of tin, about 5%, in molten copper produced an alloy called bronze that was easier to work and much harder than copper.[18] This discovery so revolutionized civilization that any culture that made widespread use of bronze to make tools and weapons became part of what archaeologists call the Bronze Age. The Bronze Age arrived in Egypt, Mesopotamia and the Indus Valley culture by around 3000 BC.[19] [20] As of 2001, the oldest tin mine found is in the Taurus Mountains in Turkey. Younger but still ancient tin mines are located in Spain, Brittany, and Great Britain.[19] European tin mining is believed to have started in Cornwall and perhaps on Dartmoor in Devon in Classical times, and a thriving tin trade developed with the civilizations of the Mediterranean.[21] [22] Securing these strategically important sites is one reason why the Romans invaded and occupied Great Britain.[19]
Ceremonial giant dirk, 1500–1300 BC.
The alchemical symbol for tin. Also used as the glyph for Jupiter.
A Bronze Age shipwreck of about 1750 BC was found at the mouth of the river Erme in Devon, with ingots of tin. A shipwreck at Uluburun, Turkey dating to 1336 BC contains a shipment of tin, perhaps originating in Afghanistan.[23] Although pure tin metal was not widely used until about 600 BC, one of the oldest tin artifacts is a ring and bottle made mostly of tin that was found in an 18th Dynasty (1580–1350 BC) tomb in Egypt, even though no tin ore reserves are known to exist in that country.[18] View from Dolcoath Mine towards Redruth, c. 1890
Tin
7
Modern times During the Middle Ages, and again in the early 19th century Cornwall was the major tin producer. This changed after large amounts of tin were found in the Bolivian tin belt and the east Asian tin belt, stretching from China through Thailand and Laos to Malaya and Indonesia. Tasmania also hosts deposits of historical importance, most notably Mount Bischoff and Renison Bell. In 1931 the tin producers founded the International Tin Committee, followed in 1956 by the International Tin Council, an institution to control the tin market. After the collapse of the market in October 1985 the price for tin nearly halved.[24] Today, the word "tin" is often improperly used as a generic term for any silvery metal that comes in sheets. Most everyday materials that are commonly called "tin", such as aluminium foil, beverage cans, corrugated building sheathing and tin cans, are actually made of steel or aluminium, although tin cans (tinned cans) do contain a thin coating of tin to inhibit rust. Likewise, so-called "tin toys" are usually made of steel, and may or may not have a coating of tin to inhibit rust. The original Ford Model T was known colloquially as the "Tin Lizzy".
Occurrence Tin is the 49th most abundant element in the Earth's crust, representing 2 ppm compared with 75 ppm for zinc, 50 ppm for copper, and 14 ppm for lead.[25] Tin does not occur naturally by itself, and must be
Crystals of cassiterite tin ore
extracted from a base compound, usually cassiterite (SnO2), the only commercially important source of tin, although small quantities of tin are recovered from complex sulfides such as stannite, cylindrite, franckeite, canfieldite, and teallite. Minerals with tin are almost always in association with granite rock, which when contain the mineral, have a 1% tin oxide content.[26]
Due to the higher specific gravity of tin and its resistance to corrosion, about 80% of mined tin is Tin output in 2005 from secondary deposits found downstream from the primary lodes. Tin is often recovered from granules washed downstream in the past and deposited in valleys or under sea. The most economical ways of mining tin are through dredging, hydraulic methods or open cast mining. Most of the world's tin is produced from placer deposits, which may contain as little as 0.015% tin. It was estimated in January 2008 that there were 6.1 million tons of economically recoverable primary
Tin
8 reserves, from a known base reserve of 11 million tons. Below are the nations with the 10 largest known reserves. Estimates of tin production have historically varied with the dynamics of economic feasibility and the development of mining technologies, but it is estimated that, at current consumption rates and technologies, the Earth will run out of tin that can be mined in 40 years.[27] However Lester Brown has suggested tin could run out within 20 years based on an extremely conservative extrapolation of 2% growth per year.[28]
Tin ore
World tin mine reserves and reserve base in tons Country
[29]
Reserves
Reserve Base
China
1,700,000
3,500,000
Malaysia
1,000,000
1,200,000
Peru
710,000
1,000,000
Indonesia
800,000
900,000
Brazil
540,000
2,500,000
Bolivia
450,000
900,000
Russia
300,000
350,000
Thailand
170,000
250,000
Australia
150,000
300,000
180,000
200,000
Other
Estimated economically recoverable [26] world tin reserves in million tons 1965
4,265
1970
3,930
1975
9,060
1980
9,100
1985
3,060
1990
7,100
2008
[30] 6,100
Secondary, or scrap, tin is also an important source of the metal and the recovery of tin through secondary production, or recycling of scrap tin, is increasing rapidly. While the United States has neither mined since 1993 nor smelted tin since 1989, it was the largest secondary producer, recycling nearly 14,000 tons in 2006.[29] New deposits are reported to be in southern Mongolia, and in 2009, new deposits of tin were discovered in Colombia, South America, by the Seminole Enterprises Group.[31] [32]
Tin
9
Production Tin is produced by reducing the ore with coal in a reverberatory furnace.
Mining and smelting In 2006, total worldwide tin mine production was 321,000 tons, and smelter production was 340,000 tons. From its production level of 186,300 tons in 1991, around where it had hovered for the previous decades, production of tin shot up 89%, to 351,800 tons in 2005. Most of the increase came from China and Indonesia, with the largest spike in 2004–2005, when it increased 23%. While in the 1970s Malaysia was the largest producer, with around a third of world production, it has steadily fallen, and now remains a major smelter and market center. In 2007, the People's Republic of China was the largest producer of tin, where the tin deposits are concentrated in the southeast Yunnan tin belt,[33] with 43% of the world's share, followed by Indonesia, with an almost equal share, and Peru at a distant third, reports the USGS.[30] The table below shows the countries with the largest mine production and the largest smelter output.[34]
Mine and smelter production (tons), 2006 Country
Mine Production
Smelter Production
China
114,300
129,400
Indonesia
117,500
80,933
Peru
38,470
40,495
Bolivia
17,669
13,500
Thailand
225
27,540
Malaysia
2,398
23,000
Belgium
0
8,000
Russia
5,000
5,500
Congo-Kinshasa ('08)
15,000
0
After the discovery of tin in what is now Bisie, North Kivu in the Democratic Republic of Congo in 2002, illegal production has increased there to around 15,000 tons.[35] This is largely fueling the ongoing and recent conflicts there, as well as affecting international markets.
Tin
10
Industry The ten largest companies produced most of world's tin in 2007. It is not clear which of these companies include tin smelted from the mine at Bisie, Congo-Kinshasa, which is controlled by a renegade militia and produces 15,000 tons. Most of the world's tin is traded on the London Metal Exchange (LME), from 8 countries, under 17 brands.[36]
Largest tin mining companies by production in tons Company
Polity
2006
2007
%Change
Yunnan Tin
China
52,339
61,129
16.7
PT Timah
Indonesia
44,689
58,325
30.5
Minsur
Peru
40,977
35,940
−12.3
Malay
China
52,339
61,129
16.7
Malaysia Smelting Corp
Malaysia
22,850
25,471
11.5
Thaisarco
Thailand
27,828
19,826
−28.8
Yunnan Chengfeng
China
21,765
18,000
−17.8
Liuzhou China Tin
China
13,499
13,193
−2.3
EM Vinto
Bolivia
11,804
9,448
−20.0
Gold Bell Group
China
4,696
8,000
70.9
Prices of tin were at $11,900 per ton as of Nov 24, 2008. Prices reached an all-time high of nearly $25,000 per ton in May 2008, largely because of the effect of the decrease of tin production from Indonesia, and have been volatile because of reliance from mining in Congo-Kinshasa.
Applications In 2006, the categories of tin use were solder (52%), tinplate (16%), chemicals (13%), brass and bronze (5.5%), glass (2%), and variety of other applications (11%)[37]
Tin
11
Metal or alloy Tin is used by itself, or in combination with other elements for a wide variety of useful alloys. Tin is most commonly alloyed with copper. Pewter is 85–99% tin; Babbitt metal has a high percentage of tin as well. Bronze is mostly copper (12% tin), while addition of phosphorus gives phosphor bronze. Bell metal is also a copper-tin alloy, containing 22% tin. Pewter plate
Tin plated metal from can
Tin bonds readily to iron, and is used for coating lead or zinc and steel to prevent corrosion. Tin-plated steel containers are widely used for food preservation, and this forms a large part of the market for metallic tin. A tinplate canister for preserving food was first manufactured in London in 1812. Speakers of British English call them "tins"; Americans call them "cans" or "tin cans". One thus-derived use of the slang term "tinnie" or "tinny" means "can of beer". The tin whistle is so called because it was first mass-produced in tin-plated steel.
Window glass is most often made via floating molten glass on top of molten tin (creating float glass) in order to make a flat surface (this is called the "Pilkington process").[38] Most metal pipes in a pipe organ are made of varying amounts of a tin/lead alloy, with 50%/50% being the most common. The amount of tin in the pipe defines the pipe's tone, since tin is the most tonally resonant of all metals. When a tin/lead alloy cools, the lead cools slightly faster and makes a mottled or spotted effect. This metal alloy is referred to as spotted metal. Tin foil was once a common wrapping material for foods and drugs; replaced in the early 20th century by the use of aluminium foil, which is now commonly referred to as tin foil. Hence one use of the slang term "tinnie" or "tinny" for a small pipe for use of a drug such as cannabis or for a can of beer. Tin becomes a superconductor below 3.72 K. In fact, tin was one of the first superconductors to be studied; the Meissner effect, one of the characteristic features of superconductors, was first discovered in superconducting tin crystals. The niobium-tin compound Nb3Sn is commercially used as wires for superconducting magnets, due to the material's high critical temperature (18 K) and critical magnetic field (25 T). A superconducting magnet weighing only a couple of kilograms is capable of producing magnetic fields comparable to a conventional electromagnet weighing tons.
Tin
12
Solder Tin has long been used as a solder in the form of an alloy with lead, tin comprising 5 to 70% w/w. Tin forms a eutectic mixture with lead containing 63% tin and 37% lead. Such solders are primarily used for solders for joining pipes or electric circuits. Since the European Union Waste Electrical and Electronic Equipment Directive (WEEED) and Restriction of Hazardous Substances Directive (RoHS) came into effect on 1 July 2006, the use of lead in such alloys has decreased. A coil of lead-free solder wire Replacing lead has many problems, including a higher melting point, and the formation of tin whiskers causing electrical problems. Replacement [39] alloys are rapidly being found.
Organotin compounds Organotin compounds have the widest range of uses of all main-group organometallic compounds, with an annual worldwide industrial production of probably exceeding 50,000 tonnes. Their major application is in the stabilization of halogenated PVC plastics, which would otherwise rapidly degrade under heat, light, and atmospheric oxygen, to give discolored, brittle products. It is believed that the tin scavenges labile chlorine ions (Cl-), which would otherwise initiate loss of HCl from the plastic material.[11] Organotin compounds have a relatively high toxicity, and for this they have been used for their biocidal effects in/as fungicides, pesticides, algacides, wood preservatives, and antifouling agents.[40] Tributyltin oxide is used as a wood preservative. Tributyltin was used as additive for ship paint to prevent growth of marine organisms on ships. The use declined after organotin compounds were recognized as persistent organic pollutants with a extremely high toxicity for some marine organisms, for example the dog whelk.[41] The EU banned the use of organotin compounds in 2003.[42] Concerns over toxicity of these compounds to marine life and their effects over the reproduction and growth of some marine species,[40] (some reports describe biological effects to marine life at a concentration of 1 nanogram per liter) have led to a worldwide ban by the International Maritime Organization. Many nations now restrict the use of organotin compounds to vessels over 25 meters long.[40] The Stille reaction couples organotin compounds with organic halides or pseudohalides.[43]
Precautions Tin plays no known natural biological role in humans, and possible health effects of tin are a subject of dispute. Tin itself is not toxic but most tin salts are. The corrosion of tin plated food cans by acidic food and beverages has caused several intoxications with soluble tin compounds. Nausea, vomiting and diarrhoea have been reported after ingesting canned food containing 200 mg/kg of tin.[44] This observation led, for example, the Food Standards Agency in the UK to propose upper limits of 200 mg/kg.[45] A study showed that 99.5% of the controlled food cans contain tin in an amount below that level.[46] Organotin compounds are very toxic. Tri-n-alkyltins are phytotoxic and, depending on the organic groups, can be powerful bactericides and fungicides. Other triorganotins are used
Tin
13 as miticides and acaricides.
See also • • • • • • •
Stannary Tinning Cassiterides (the mythical Tin Islands) Tin pest Whisker (metallurgy) (tin whiskers) Terne Tin mining in Britain
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14 [26] " Tin: From Ore to Ingot (http:/ / www. itri. co. uk/ pooled/ articles/ BF_TECHART/ view. asp?Q=BF_TECHART_230527)". International Tin Research Institute. 1991. . Retrieved 2009-03-21. [27] "How Long Will it Last?". New Scientist 194 (2605): 38–39. May 26, 2007. ISSN 4079 0262 4079 (http:/ / worldcat. org/ issn/ 0262). [28] Brown, Lester (2006). Plan B 2.0. New York: W.W. Norton. pp. 109. ISBN 978-0393328318. [29] Carlin, Jr., James F.. " Minerals Yearbook 2006: Tin (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ tin/ myb1-2006-tin. pdf)" (PDF). United States Geological Survey. . Retrieved 2008-11-23. [30] Carlin, Jr., James F.. " Mineral Commodity Summary 2008: Tin (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ tin/ mcs-2008-tin. pdf)" (PDF). United States Geological Survey. . [31] " Seminole Group Colombia Discovers High Grade Tin Ore in the Amazon Jungle (http:/ / www. 1888pressrelease. com/ seminole-group-colombia-discovers-high-grade-tin-ore-in-the-pr-100235. html)". 1888 PressRelease. . Retrieved 2009-07-28. [32] " Seminole Enterprises Group Discovers High Grade Tin Ore In The Amazons Of Colombia (http:/ / www. prlog. org/ 10175604-seminole-enterprises-group-discovers-high-grade-tin-ore-in-the-amazons-of-colombia. html)". PRLog Free Press Release. . Retrieved 2009-07-28. [33] Shiyu, Yang (1991). "Classification and type association of tin deposits in Southeast Yunnan Tin Belt". Chinese Journal of Geochemistry 10 (1): 21–35. doi: 10.1007/BF02843295 (http:/ / dx. doi. org/ 10. 1007/ BF02843295). [34] Estimates vary between USGS and The British Geological Survey. The latter was chosen because it indicates that the most recent statistics are not estimates, and estimates match more closely with other estimates found for Congo-Kinshasa. [35] " The Spoils: Congo's Riches, Looted by Renegade Troops (http:/ / www. nytimes. com/ 2008/ 11/ 16/ world/ africa/ 16congo. html?ref=africa)". New York Times. November 15, 2008. . [36] " International Tin Research Institute. LME Tin Brands (http:/ / www. itri. co. uk/ pooled/ articles/ BF_TECHART/ view. asp?Q=BF_TECHART_303032)". . Retrieved 2009-05-05. [37] " ITRI. Tin Use Survey 2007 (http:/ / www. itri. co. uk/ pooled/ articles/ BF_TECHART/ view. asp?Q=BF_TECHART_297350)". ITRI. . Retrieved 2008-11-21. [38] Pilkington, L. A. B.. " Review Lecture. The Float Glass Process (http:/ / www. jstor. org/ stable/ 2416528)". Proceedings of the Royal Society of London. Series A, Mathematical and Physical Sciences 314 (1516): 1–25. . [39] Black, Harvey. (2005). " Getting the Lead out of Electronics (http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?artid=1281311)". Environmental Health Perspectives 113 (10). . [40] . p. 345. ISBN 0716748789. [41] Eisler, Ronald. " Tin Hazards To Fish, Wildlife, and Invertebrates: A Synoptic Review (http:/ / www. dtic. mil/ cgi-bin/ GetTRDoc?AD=ADA322822& Location=U2& doc=GetTRDoc. pdf)" (PDF). U.S. Fish and Wildlife Service Patuxent Wildlife Research Center. . [42] " Regulation (EC) No 782/2003 of the European Parlament and of the Council of 14 April 2003 on the prohibition of organotin compounds on ships (http:/ / eur-lex. europa. eu/ LexUriServ/ LexUriServ. do?uri=OJ:L:2003:115:0001:0011:EN:PDF)". . Retrieved 2009-05-05. [43] Farina, Vittorio; Krishnamurthy, Venkat; Scott, William J. (1997). "The Stille Reaction". Organic Reactions. doi: 10.1002/0471264180.or050.01 (http:/ / dx. doi. org/ 10. 1002/ 0471264180. or050. 01). [44] Blunden, Steve; Wallace, Tony (2003). "Tin in canned food: a review and understanding of occurrence and effect". Food and Chemical Toxicology 41 (12): 1651–1662. doi: 10.1016/S0278-6915(03)00217-5 (http:/ / dx. doi. org/ 10. 1016/ S0278-6915(03)00217-5). [45] " Eat well, be well — Tin (http:/ / www. eatwell. gov. uk/ healthissues/ factsbehindissues/ tins/ )". Food Standards Agency. . Retrieved 2009-04-16. [46] " Tin in canned fruit and vegetables (Number 29/02) (http:/ / www. food. gov. uk/ multimedia/ pdfs/ fsis2902tin. pdf)" (PDF). Food Standards Agency. 2002-08-22. . Retrieved 2009-04-16.
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Bibliography • CRC contributors (2006). David R. Lide (editor). ed. Handbook of Chemistry and Physics (87th ed.). Boca Raton, Florida: CRC Press, Taylor & Francis Group. ISBN 0-8493-0487-3. • Emsley, John (2001). "Tin". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. pp. 445–450. ISBN 0198503407. • Stwertka, Albert (1998). "Tin". Guide to the Elements (Revised ed.). Oxford University Press. ISBN 0-19-508083-1. • Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. ISBN 0-7506-3365-4. • MacIntosh, Robert M. (1968). "Tin". in Clifford A. Hampel (editor). The Encyclopedia of the Chemical Elements. New York: Reinhold Book Corporation. pp. 722–732. LCCN 68-29938. • Heiserman, David L. (1992). "Element 50: Tin". Exploring Chemical Elements and their Compounds. New York: TAB Books. ISBN 0-8306-3018-X.
External links • WebElements.com – Tin (http:/ / www. webelements. com/ webelements/ elements/ text/ Sn/ index. html) • Theodore Gray's Wooden Periodic Table Table (http:/ / www. theodoregray. com/ PeriodicTable/ Elements/ 050/ index. s7. html): Tin samples and castings • Base Metals: Tin (http:/ / www. basemetals. com/ html/ sninfo. htm) • Comprehensive Data on Tin (http:/ / www. mrteverett. com/ Chemistry/ pdictable/ q_elements. asp?Symbol=Sn)
Article Sources and Contributors
Article Sources and Contributors Tin Source: http://en.wikipedia.org/w/index.php?oldid=305322506 Contributors: 21655, A2Kafir, Aadal, Abdclmop, Ace Telephone, Acroterion, Adashiel, Addshore, Adresia, Ahoerstemeier, Ahwmay, AirdishStraus, Aitias, Ale jrb, AlexiusHoratius, [email protected], AliveFreeHappy, Amire80, AndonicO, Andromedahpg, Ann Stouter, Antewolf, Anthony Appleyard, Anwar saadat, Arakunem, Archimerged, Ardric47, Arkuat, Astral highway, Auric, Avono, Axiosaurus, B. van der Wee, Beetstra, Beland, Benbest, Bencherlite, Bennetto, Billy ruth, Blanchardb, BlueEarth, Bobianite, Bobo192, Bobobobo94, Bongwarrior, Brianhe, Bryan Derksen, Bsimmons666, Butros, C.Fred, CIreland, CYD, Can't sleep, clown will eat me, CanadianLinuxUser, CanisRufus, Capricorn42, CaptainVindaloo, CardinalDan, CatherineMunro, CharlesC, Charlotte96, Chris 73, Chris the speller, Christian List, Chun-hian, Cimon Avaro, Cobi, Cometstyles, Conversion script, Cool3, Coppelia, Coppertwig, Cryptic C62, Cybercobra, CzarB, DVD R W, Dan100, Danny, Darkwraith, Darrien, Darth Panda, Dave Farquhar, Davgold, David Latapie, Decltype, Dekker, Delta G, Dennis Brown, DerHexer, Dewnans, Dgg2, Dino, Diplogkop, Discospinster, Dlae, Docboat, Donarreiskoffer, DoubleBlue, Dougofborg, Doulos Christos, Dreamingkat, Drestros power, Drini, Duffman, Dughnjghfsg, DuncanHill, Dyvroeth, E rulez, Eclipsenow, Edgar181, El C, Emperorbma, Enviroboy, Epbr123, Eric-Wester, Erik Zachte, Ernst de 2nd q.u., Escape Orbit, Euchiasmus, Everyking, Excirial, F, FJPB, Femto, Firefoxman, Flaminghomeryto, FlyingToaster, Frankenpuppy, Frankenstud, FrederikRamm, Frehley, Froboz23, Fvw, Fyyer, GHe, Gene Nygaard, Giftlite, Gilliam, GraemeL, Grendelkhan, Greyhood, Gromlakh, Gurch, Guy1000, Gwernol, Haade, Hadrianheugh, HenryLi, Herbee, HereToHelp, Heron, Hgrosser, Hiddekel, Honeycake, I love scarecrows, ILike2BeAnonymous, Icairns, IdLoveOne, Ideyal, Indosauros, It Is Me Here, Itub, Ixfd64, J.delanoy, JIP, Jaan513, James A. Stewart, Jamyskis, Janfri, Jaraalbe, Jauerback, Jeberle, JeffAMcGee, JerrySteal, Jerzy, Jh51681, Jhfjashfjhafhiusdh, Jhinman, Jivnak, Joanjoc, John, JonC, Jose77, Josh Grosse, JoshuaZ, Kairos, Karlhahn, Katieh5584, Kember, Kessler, Klippa, Ktsquare, Kurykh, LastChanceToBe, Laurinavicius, Leal Nightrunner, Leonard G., Lethalgeek, Lightmouse, Listering, Lost tourist, LovesMacs, Luckas Blade, LuigiManiac, MC10, MSmolik, Mac, Macintosh User, Mackeroj16, Magnus Bakken, Makemi, Mapleleafsrules, Martpol, Materialscientist, Mav, Maxamegalon2000, Maximus Rex, Mboverload, McTrixie, Mejor Los Indios, Melos Antropon, Mervyn, Metsfanmax, Mgimpel, Mightypile, Mike Young, Mike998877665544332211, Mikeblas, Minerallad, Minesweeper, Motley Crue Rocks, Moverton, MrDolomite, Müslimix, NHRHS2010, NawlinWiki, Neddyseagoon, Neon white, Nergaal, Neurolysis, Nick C, Nihiltres, NittyG, Nivix, Noe, Novalis, Nwbeeson, Oda Mari, Old Moonraker, OllieFury, Onceler, Onevalefan, Oroso, Ossmann, Oxymoron83, PMDrive1061, Paraballo, Pb30, Pearle, Pevernagie, Pharillon, PigFlu Oink, Plantsurfer, Plasticup, PlatinumX, Pleasantville, Pollinator, Poolkris, Pparazorback, Psyche825, Quintote, RDT2, RTC, Rebrane, RedWolf, Redmailbox1, Reedy, Remember, Res2216firestar, Retard107, RexNL, Rgoodermote, RichG, Rifleman 82, Rjwilmsi, RobertG, Robin Patterson, Rock hopper55, Rocket71048576, Romanm, Roni2204, Rossnorman, SGBailey, Sachin318, Saperaud, Sbharris, Sceptre, SchfiftyThree, Schneelocke, Selkem, Sengkang, Senor Kittykins, Seraphim, Serlin, Sfmammamia, Shaddack, Shappy, ShaunMacPherson, Shoeofdeath, Shoessss, Silentlight, Sillybilly, Silverbackmarlin, Sionus, Skizzik, SkyWalker, Sl, Smalljim, Snigbrook, Snowolf, Special-T, SpeedyGonsales, SpuriousQ, Squids and Chips, Stephenb, Stone, Storm Rider, Suisui, Svante, Swansrule, T g7, Tagishsimon, Tangent747, Tbone55, Tetracube, The Photon, The Random Editor, TheDJ, TheMidnighters, Thedjatclubrock, Theseeker4, Theunixgeek, Thienghiem, Thricecube, Tickopa, Tide rolls, Time3000, Tobyc75, Tomatensaft, Tonkerbelle91, Triwbe, Tropylium, Turnstep, Uppermile, V8rik, Versus22, Violinchick1995, Vlad4599, Vsmith, Vssun, Wadester16, WarthogDemon, Warut, Watch37264, Werdan7, Wiki alf, Wildwolfboy11, Will Beback, Willtron, Wimt, Woland37, Worrydream, Wrenchelle, Wtmitchell, Xy7, Yath, Yekrats, Yyy, Zero sharp, Ziusudra, 762 anonymous edits
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17
Antimony
1
Antimony tin ← antimony → telluriumAs ↑ Sb ↓ Bi
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
51Sb Periodic table
Appearance silvery lustrous gray
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
Antimony
2 at T/K
807
876
1011
1219
1491
1858
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1594.9 kJ·mol−1 3rd: 2440 kJ·mol−1Atomic radiusCovalent radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusMohs hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of antimony iso
N.A.
half-life
121
57.36%
121
123
42.64%
123
125
syn
2.7582 y
Sb Sb Sb
DM
DE (MeV)
DP
Sb is stable with 70 neutron Sb is stable with 72 neutron β−
0.767
125
Te
antimony, Sb, 51 metalloid15, 5, p121.760(1) g·mol−1 [Kr] 4d10 5s2 5p3 2, 8, 18, 18, 5 (Image) solid 6.697 g·cm−3 6.53 g·cm−3 903.78 K,630.63 °C,1167.13 °F 1860 K,1587 °C,2889 °F 19.79 kJ·mol−1 193.43 kJ·mol−1 (25 °C) 25.23 J·mol−1·K−15, 3, -3 2.05 (Pauling scale) 1st: 834 kJ·mol−1140 pm139±5 pm rhombohedral diamagnetic[1] (20 °C) 417 nΩ·m (300 K) 24.4 W·m−1·K−1 (25 °C) 11 µm·m−1·K−1 (20 °C) 3420 m/s 55 GPa 20 GPa 42 GPa 3.0 294 MPa 7440-36-0 Antimony (pronounced /ænˈtɪmənɪ/ (UK) or English pronunciation: /ˈæntɨmoʊni/ (US)) is a chemical element with the symbol Sb (Latin: stibium, meaning "mark") and atomic number 51. A metalloid, antimony has four allotropic forms. The stable form of antimony is a blue-white metalloid. Yellow and black antimony are unstable non-metals. Antimony is used in flame-proofing, paints, ceramics, enamels, a wide variety of alloys, electronics, and rubber.
Properties Antimony in its elemental form is a silvery white, brittle, fusible, crystalline solid that exhibits poor electrical and heat conductivity properties and vaporizes at low temperatures. A metalloid, antimony resembles a metal in its appearance and in many of its physical properties, but does not chemically react as a metal. It is also attacked by or reactive with oxidizing acids and halogens. Antimony and some of its alloys are unusual in that they expand on cooling. Antimony is geochemically categorized as a chalcophile, occurring with sulfur and the heavy metals lead, copper, and silver. The abundance of antimony in the Earth's crust is estimated at 0.2 to 0.5 parts per million.[2]
Antimony
Applications Antimony is increasingly being used in the semiconductor industry in the production of diodes, infrared detectors, and Hall-effect devices. As an alloy, this metalloid greatly increases lead's hardness and mechanical strength. The most important use of antimony is as a hardener in lead for storage batteries. Uses include: • • • • • • • • • • • •
Batteries antifriction alloys small arms, buckshot, and tracer ammunition cable sheathing matches medicines, antiprotozoan drugs plumbing soldering - some "lead-free" solders contain 5% Sb main and big-end bearings in internal combustion engines (as alloy) used in the past to treat Schistosomiasis; today Praziquantel is universally used used in type metal, e.g. for linotype printing machines used in pewter
Antimony compounds in the form of oxides, sulfides, sodium antimonate, and antimony trichloride are used in the making of flame-proofing compounds, ceramic enamels, glass, paints, and pottery. Antimony trioxide is the most important of the antimony compounds and is primarily used in flame-retardant formulations. These flame-retardant applications include such markets as children's clothing, toys, aircraft and automobile seat covers. It is also used in the fiberglass composites industry as an additive to polyester resins for such items as light aircraft engine covers. The resin will burn while a flame is held to it but will extinguish itself as soon as the flame is removed. Antimony sulfide is also one of the ingredients of safety matches. In the 1950s, tiny beads of a lead-antimony alloy were used for the emitters and collectors of NPN alloy junction transistors. The natural sulfide of antimony, stibnite, was known and used in Biblical times, as medicine and in Islamic/Pre-Islamic times as a cosmetic. The Sunan Abi Dawood reports, “Muhammad said: 'Among the best types of collyrium use is antimony (ithmid) for it clears the vision and makes the hair sprout.'”[3] Stibnite is still used in some developing countries as medicine. Antimony has been used for the treatment of schistosomiasis. Antimony attaches itself to sulfur atoms in certain enzymes which are used by both the parasite and human host. Small doses can kill the parasite without causing damage to the patient. Antimony and its compounds are used in several veterinary preparations like Anthiomaline or Lithium antimony thiomalate, which is used as a skin conditioner in ruminants. Antimony has a nourishing or conditioning effect on keratinized tissues, at least in animals. Tartar emetic is another antimony preparation which is used as an anti-schistosomal drug. Treatments chiefly involving antimony have been called antimonials. Antimony-based drugs such as allopurinol and Meglumine, are also considered the drugs of choice for the treatment of leishmaniasis in domestic animals. Unfortunately, as well as having low therapeutic indices, the drugs are poor at penetrating the bone marrow, where some of the Leishmania amastigotes reside, and so cure of the disease - especially the visceral form - is very difficult.
3
Antimony
4
A coin made of antimony was issued in the Keichow Province of China in 1931. The coins were not popular, being too soft and they wore quickly when in circulation. After the first issue no others were produced.[4]
Etymology The ancient words for antimony mostly have, as their chief meaning, kohl, the sulfide of antimony. Pliny the Elder, however, distinguishes between male and female forms of antimony; his male form is probably the sulfide, the female form, which is superior, heavier, and less friable, is probably native metallic antimony.[5] The Egyptians called antimony mśdmt; in hieroglyphics, the vowels are uncertain, but there is an Arabic tradition that the word is mesdemet.[6] The Greek word, stimmi, is probably a loan word from Arabic or Egyptian, and is used by the Attic tragic poets of the 5th century BC; later Greeks also used stibi, as did Celsus and Pliny, writing in Latin, in the first century AD. Pliny also gives the names stimi [sic], larbaris, alabaster, and the "very common" platyophthalmos, "wide-eye" (from the effect of the cosmetic). Later Latin authors adapted the word to Latin as stibium. The Arabic word for the substance, as opposed to the cosmetic, can appear as ithmid, athmoud, othmod, or uthmod. Littré suggests the first form, which is the earliest, derives from stimmida, (one) accusative for stimmi.[7] The use of Sb as the standard chemical symbol for antimony is due to the 18th century chemical pioneer, Jöns Jakob Berzelius, who used this abbreviation of the name stibium. The medieval Latin form, from which the modern languages and late Byzantine Greek, take their names, is antimonium. The origin of this is uncertain; all suggestions have some difficulty either of form or interpretation. The popular etymology, from anti-monachos or French antimoine, still has adherents; this would mean "monk-killer", and is explained by many early alchemists being monks, and antimony being poisonous.[8] So does the hypothetical Greek word antimonos, "against one", explained as "not found as metal", or "not found unalloyed".[9] Lippmann conjectured a Greek word, anthemonion, which would mean "floret", and he cites several examples of related Greek words (but not that one) which describe chemical or biological efflorescence.[10] The early uses of antimonium include the translations, in 1050-1100, by Constantine the African of Arabic medical treatises.[11] Several authorities believe that antimonium is a scribal corruption of some Arabic form; Meyerhof derives it from ithmid;[12] other possibilities include Athimar, the Arabic name of the metal, and a hypothetical *as-stimmi, derived from or parallel to the Greek.[13]
Antimony
5
History
One of the alchemical symbols for antimony.
Antimony's sulfide compound, antimony (III) trisulfide, Sb2S3 was recognized in antiquity, at least as early as 3000 BC. Pastes of Sb2S3 powder in fat[14] or in other materials have been used since that date as eye cosmetics in the Middle East and farther afield; in this use, Sb2S3 is called kohl. It was used to darken the brows and lashes, or to draw a line around the perimeter of the eye.
An artifact made of antimony dating to about 3000 BC was found at Tello, Chaldea (part of present-day Iraq), and a copper object plated with antimony dating between 2500 BC and 2200 BC has been found in Egypt.[15] There is some uncertainty as to the description of the artifact from Tello. Although it is sometimes reported to be a vase, a recent detailed discussion of reports it to be rather a fragment of indeterminate purpose.[16] The first European description of a procedure for isolating antimony is in the book De la pirotechnia of 1540 by Vannoccio Biringuccio, written in Italian. This book precedes the more famous 1556 book in Latin by Agricola, De re metallica, even though Agricola has been often incorrectly credited with the discovery of metallic Native massive antimony with antimony. A text describing the preparation of metallic oxidation products antimony that was published in Germany in 1604 purported to date from the early fifteenth century, and if authentic it would predate Biringuccio. The book, in German, was the Triumph Wagen Antimonii (Triumphal Chariot of Antimony), and its putative author was a certain Benedictine monk, writing under the name Basilius Valentinus. Already in 1710 Wilhelm Gottlob Freiherr von Leibniz, after careful inquiry, concluded that the work was spurious, that there was no monk named Basilius Valentinus, and the book's author was its ostensible editor, Johann Thölde (ca. 1565-ca. 1624). There is now agreement among professional historians that the Triumph Wagen was written after the middle of the sixteenth century and that Thölde was likely its author.[17] An English translation of the Triumph Wagen appeared in English in 1660, under the title The Triumphant Chariot of Antimony. The work remains of great interest, chiefly because it documents how followers of the renegade German physician, Philippus Theophrastus Paracelsus von Hohenheim (of whom Thölde was one), came to associate the practice of alchemy with the preparation of chemical medicines. According to the traditional history of Middle Eastern alchemy, pure antimony was well known to Geber, sometimes called "the Father of Chemistry", in the 8th century. Here there is still an open controversy: Marcellin Berthelot, who translated a number of Geber's books, stated that antimony is never mentioned in them, but other authors claim that Berthelot translated only some of the less important books, while the more interesting ones (some of which might describe antimony) are not yet translated, and their content is completely
Antimony
6
unknown. The first natural occurrence of pure antimony ('native antimony') in the Earth's crust was described by the Swedish scientist and local mine district engineer Anton von Swab in 1783. The type-sample was collected from the Sala Silver Mine in the Bergslagen mining district of south central Sweden.
Production Even though this element is not abundant, it is found in over 100 mineral species. Antimony is sometimes found native, but more frequently it is found in the sulfide stibnite (Sb2S3) which is the predominant ore mineral. Commercial forms of antimony are generally ingots, broken pieces, granules, and cast cake. Other forms are powder, shot, and single crystals.
Antimony output in 2005
In 2005, China was the top producer of antimony with about 84% world share followed at a distance by South Africa, Bolivia and Tajikistan, reports the British Geological Survey.
World production trend of antimony
Country
Tonnes
% of total
People's Republic of China
126,000
84.0
South Africa
6,000
4.0
Bolivia
5,225
3.5
Tajikistan
4,073
2.7
Russia
3,000
2.0
144,298
96.2
Top 5 Total world
150,000
100.0
Chiffres de 2003, métal contenue dans les minerais et concentrés, source: L'état du monde 2005
The largest mine in China is Xikuangshan mine in Hunan Province.
Antimony
Precautions Antimony and many of its compounds are toxic. Clinically, antimony poisoning is very similar to arsenic poisoning. In small doses, antimony causes headache, dizziness, and depression. Larger doses cause violent and frequent vomiting, and will lead to death in a few days. Antimony leaches from polyethylene terephthalate (PET) bottles into bottled water, but at levels below drinking water guidelines.[18] [19] The guidelines are: • World Health Organization: 20 µg/L • Japan: 15 µg/L[20] • United States Environmental Protection Agency, Health Canada and the Ontario Ministry of Environment: 6 µg/L • German Federal Ministry of Environment: 5 µg/L[21] The acidic nature of the drink is sufficient to dissolve small amounts of antimony trioxide contained in the packaging of the drink; modern manufacturing methods prevent this occurrence. The longer the beverage has been bottled and the higher the temperature, the more antimony is leached.[22]
Compounds Important compounds of antimony include: • • • • •
Antimony pentafluoride SbF5 Antimony trioxide Sb2O3 Stibine (antimony trihydride SbH3) Indium antimonide (InSb) Fluoroantimonic acid (HSbF6)
See also • Antimonial • Phase change memory • Naturalis Historia • Pliny the Elder
Bibliography • W. F. Albright "Notes on Egypto-Semitic Etymology. II" [23], The American Journal of Semitic Languages and Literatures, Vol. 34, No. 4. (Jul., 1918), pp. 215–255 (p.230) • Endlich, F.M. "On Some Interesting Derivations of Mineral Names" [24], The American Naturalist, Vol. 22, No. 253. (Jan., 1888), pp. 21–32 (p.28) • Kirk-Othmer Encyclopedia of Chemical Technology, 5th ed. 2004. Entry for antimony. • Edmund Oscar von Lippmann (1919) Entstehung und Ausbreitung der Alchemie, teil 1. Berlin: Julius Springer. In German. • Moorey, PRS. (1994) Ancient Mesopotamian Materials and Industries: the Archaeological Evidence. New York: Clarendon Press. • Priesner, Claus and Figala, Karin, eds. (1998) Alchemie. Lexikon einer hermetischen Wissenschaft. München: C.H. Beck. 412 p. In German.
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Antimony
8
• Sarton, George. (1935) Review [25] of Al-morchid fi'l-kohhl, ou Le guide d'oculistique, translated by Max Meyerhof. Isis (1935), 22(2):539-542 In French. • Shotyk, W; Krachler, M; Chen, B (Feb 2006). "Contamination of Canadian and European bottled waters with antimony from PET containers.". Journal of environmental monitoring : JEM 8 (2): 288–92. doi:10.1039/b517844b [26]. ISSN 1464-0325 [27]. PMID 16470261. • Public Health Statement for Antimony [28] • Wakayama, Hiroshi, "Revision of Drinking Water Standards in Japan" Health, Labor and Welfare (Japan), 2003
[29]
, Ministry of
External links • National Pollutant Inventory - Antimony and compounds • WebElements.com – Antimony [31]
[30]
pnb:ینومیٹنیا
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] " Antimony Statistics and Information (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ antimony/ )". United States Geological Survey. 2009-01-31. . Retrieved 2009-04-15. [3] Sunan Abu-Dawud (Ahmad Hasan translation). Book 32, Number 4050 (http:/ / www. muslimaccess. com/ sunnah/ hadeeth/ abudawud/ 032. html). . [4] " Metals Used in Coins and Medals (http:/ / www. tclayton. demon. co. uk/ metal. html)". Tclayton.demon.co.uk. . Retrieved 2008-09-12. [5] Pliny, Natural history, 33.33; W.H.S. Jones, the Loeb Classical Library translator, supplies a note suggesting the identifications. [6] Albright, p.230; Sarton p.541, quotes Meyerhof, the translator of the book he is reviewing. [7] LSJ, s.v., vocalisation, spelling, and declension vary; Endlich, p.28; Celsus, 6.6.6 ff; Pliny Natural History 33.33; Lewis and Short: Latin Dictionary. OED, s. "antimony". [8] The use of a symbol resembling an upside down "female" symbol for antimony could also hint at a satirical pun in this origin [9] See, for example, Diana Fernando, Alchemy : an illustrated A to Z (1998) and Kirk-Othmer (below) respectively. Fernando even derives it from the story of how "Basil Valentine" and his fellow monastic alchemists poisoned themselves by working with antimony; antimonium is found two centuries before his time. "Popular etymology" from OED; as for antimonos, the pure negative would be more naturally expressed by a"not". [10] Lippman, p.643-5 [11] Lippman, p.642, writing in 1919, says "zuerst". [12] Meyerhof as quoted in Sarton, p.541, asserts that ithmid or athmoud became corrupted in the medieval "traductions barbaro-latines".; the OED asserts that some Arabic form is the origin, and if ithmid is the root, posits athimodium, atimodium, atimonium, as intermediate forms. [13] Endlich, p.28; one of the advantages of as-stimmi would be that it has a whole syllable in common with antimonium. [14] Priesner and Figala [15] Kirk-Othmer, entry "Antimony" [16] The fragment was presented in a lecture in 1892. One contemporary commented, "we only know of antimony at the present day as a highly brittle and crystalline metal, which could hardly be fashioned into a useful vase, and therefore this remarkable 'find' must represent the lost art of rendering antimony malleable." Moorey 1994:241 [17] E.g., Claus Priesner and Karin Figala, eds. (1998), Alchemie: Lexikon einer hermetischen Wissenschaft (Munich: Beck), s.v. "Basilius Valentinus." Harold Jantz was perhaps the only modern scholar to deny Thölde's authorship, but he too agrees that the work dates from after 1550: see his catalogue of German Baroque literature, available online at (http:/ / microformguides. gale. com/ Data/ Download/ 2025000R. pdf). [18] Shotyk, W; Krachler, M; Chen, B (Feb 2006). "Contamination of Canadian and European bottled waters with antimony from PET containers.". Journal of environmental monitoring : JEM 8 (2): 288–92. doi:
Antimony 10.1039/b517844b (http:/ / dx. doi. org/ 10. 1039/ b517844b). ISSN 1464-0325 (http:/ / worldcat. org/ issn/ 1464-0325). PMID 16470261. [19] " London Free Press: (http:/ / www. lfpress. com/ cgi-bin/ publish. cgi?p=120232& x)". Lfpress.com. . Retrieved 2008-09-12. [20] H. Wakayama, Table 2, p. 84 [21] Shotyk et al., 2006 [22] Westerhoff, P; Prapaipong P, Shock E, Hillaireau A (February 2008). "Antimony leaching from polyethylene terephthalate (PET) plastic used for bottled drinking water". Water Research 42 (3): 551–556. doi: 10.1016/j.watres.2007.07.048 (http:/ / dx. doi. org/ 10. 1016/ j. watres. 2007. 07. 048). PMID 17707454. [23] http:/ / links. jstor. org/ sici?sici=1062-0516%28191807%2934%3A4%3C215%3ANOEEI%3E2. 0. CO%3B2-J [24] http:/ / links. jstor. org/ sici?sici=0003-0147%28188801%2922%3A253%3C21%3AOSIDOM%3E2. 0. CO%3B2-W [25] http:/ / links. jstor. org/ sici?sici=0021-1753%28193502%2922%3A2%3C539%3A%28FOLGD%3E2. 0. CO%3B2-L [26] [27] [28] [29] [30] [31]
http:/ / dx. doi. org/ 10. 1039%2Fb517844b http:/ / worldcat. org/ issn/ 1464-0325 http:/ / www. atsdr. cdc. gov/ toxprofiles/ phs23. html http:/ / www. nilim. go. jp/ lab/ bcg/ siryou/ tnn/ tnn0264pdf/ ks0264011. pdf http:/ / www. npi. gov. au/ database/ substance-info/ profiles/ 10. html http:/ / www. webelements. com/ webelements/ elements/ text/ Sb/ index. html
9
Article Sources and Contributors
Article Sources and Contributors Antimony Source: http://en.wikipedia.org/w/index.php?oldid=308284953 Contributors: A new name 2008, A2Kafir, AWP1012933, Aerion, Ahoerstemeier, Anagnorisis, Anclation, Andre Engels, Andres, Antandrus, Anthony, Anwar saadat, Aramgutang, Archimerged, Ashumaloz, Aussie Alchemist, Axl, Badagnani, Beetstra, Benbest, Bentu, Bezking, BillC, Blanchardb, BlueEarth, Borislav.dopudja, Branddobbe, Branden, Brian0918, Brice one, Bryan Derksen, CYD, Cacahueten, Carnildo, Cerebrith, Charles Matthews, Chemkid1, Cjmnyc, Clappingsimon, CombatCraig, Conversion script, Corpx, Cpt Adham, Crag, Cybercobra, Daevatgl, Daniel2500, Danny B-), Darrien, Davidryan168, Dejvid, Delta G, Dmichaelharris, Dogcow, DrFO.Jr.Tn, Drc79, Dreish, DrewT2, Dtgm, EPO, Earin, Edgar181, Edward, El C, Emperorbma, Encyclopedia77, Epbr123, Erik Zachte, EthanL, Evda, Everyking, Femto, Fiveless, FlyingPenguins, FocalPoint, FullMetalJacket, GRAHAMUK, Gaius Cornelius, Gene Nygaard, Goudzovski, GregorB, Grendelkhan, Greyhood, Hairy Dude, Hannibal, Harish2k1vet, Haydonnoel, Hdante, Hellbus, HenryLi, Henrygb, HeroGiant, Heron, Hhaithait, Hurmata, Icairns, Ino5hiro, Ioeth, Iorsh, Itub, Ixfd64, J.delanoy, J8079s, Jagged 85, James A. Stewart, Jaraalbe, Jedager1243, Jerzy, Jmundo, Joanjoc, John, John254, Jorfer, Jose77, Joshuaali, Keenan Pepper, King of Hearts, Kjkolb, Koavf, Koyaanis Qatsi, Kukini, Kurykh, Kwamikagami, LA2, Leyo, Lon of Oakdale, Looxix, Lusanaherandraton, Lvl, M.e, Mailer diablo, Manasij, Mandel, Materialscientist, Mattbaileyisneat, Mav, Mboverload, Megan1967, Mentifisto, Mervyn, Miltonhowe, Minesweeper, Mirv, Mmarko, Mydoctor93, Nabokov, Namangwari, NawlinWiki, Nergaal, NewEnglandYankee, Nickptar, Nihiltres, Nlu, Nunquam Dormio, Oliphaunt, Otterwiki, PP Jewel, Palica, Parsival, Paul Stansifer, Pax:Vobiscum, Pcb21, PeterJeremy, Peterlewis, Piano non troppo, PlatinumX, Pmanderson, Poolkris, Power.corrupts, Pras, Quadell, R6144, RTC, Remember, Reuvenk, RexNL, Reza kalani, Rjanag, Rjboyer, Rjstott, Robert McClenon, Roberta F., Romanm, Rosser1954, Rursus, Sambc, Sanctu, Saperaud, SassyLee, Sceder, Schneelocke, Schzmo, Scorpion451, Scottydude, Scyrene, Sengkang, Shadowin, Shafik, Shirt58, Skatebiker, Sl, Soliloquial, Spencer, SpuriousQ, Squids and Chips, Srleffler, Srnec, SteinbDJ, Stifynsemons, Stillnotelf, Stone, Tetracube, Verne Equinox, Vsmith, WackyBoots, Walton One, Warut, Watch37264, Wereon, Wik, Wiki alf, Wolf364, Xuanji, YellowMonkey, Yyy, Zoicon5, 263 anonymous edits
Image Sources, Licenses and Contributors file:rhombohedral.svg Source: http://en.wikipedia.org/w/index.php?title=File:Rhombohedral.svg License: unknown Contributors: User:Stannered file:Electron shell 051 Antimony.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_051_Antimony.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 file:Antimon Barren von 16 kg.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Antimon_Barren_von_16_kg.jpg License: GNU Free Documentation License Contributors: Cdang, Saperaud Image:Antimon.PNG Source: http://en.wikipedia.org/w/index.php?title=File:Antimon.PNG License: Public Domain Contributors: Ondřej Mangl Image:antimony-symbol.svg Source: http://en.wikipedia.org/w/index.php?title=File:Antimony-symbol.svg License: unknown Contributors: User:Reuvenk Image:Antimony massive.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Antimony_massive.jpg License: unknown Contributors: User:Aramgutang Image:Antimony (mined)2.PNG Source: http://en.wikipedia.org/w/index.php?title=File:Antimony_(mined)2.PNG License: Public Domain Contributors: BloodIce, Mdd Image:Antimony - world production trend.svg Source: http://en.wikipedia.org/w/index.php?title=File:Antimony_-_world_production_trend.svg License: Public Domain Contributors: User:Leyo Image:Flag of the People's Republic of China.svg Source: http://en.wikipedia.org/w/index.php?title=File:Flag_of_the_People's_Republic_of_China.svg License: Public Domain Contributors: User:Denelson83, User:SKopp, User:Shizhao, User:Zscout370 Image:Flag of South Africa.svg Source: http://en.wikipedia.org/w/index.php?title=File:Flag_of_South_Africa.svg License: Public Domain Contributors: User:SKopp Image:Flag of Bolivia.svg Source: http://en.wikipedia.org/w/index.php?title=File:Flag_of_Bolivia.svg License: Public Domain Contributors: User:SKopp Image:Flag of Tajikistan.svg Source: http://en.wikipedia.org/w/index.php?title=File:Flag_of_Tajikistan.svg License: Public Domain Contributors: Alex Spade, Apatomerus, EugeneZelenko, Homo lupus, Klemen Kocjancic, Nameneko, Neq00, Nightstallion, 1 anonymous edits Image:Flag of Russia.svg Source: http://en.wikipedia.org/w/index.php?title=File:Flag_of_Russia.svg License: Public Domain Contributors: AndriusG, Davepape, Dmitry Strotsev, Enbéká, Fred J, Gleb Borisov, Herbythyme, Homo lupus, Kiensvay, Klemen Kocjancic, Kwj2772, Mattes, Maximaximax, Miyokan, Nightstallion, Ondřej Žváček, Pianist, Pumbaa80, Putnik, R-41, Radziun, Rainman, Reisio, Rfc1394, Rkt2312, Sasa Stefanovic, SeNeKa, SkyBon, Srtxg, Stianbh, Westermarck, Wikiborg, Winterheart, Zscout370, Zyido, ОйЛ, 34 anonymous edits
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10
Tellurium
1
Tellurium 52
antimony ← tellurium → iodine
Se ↑
Te ↓
Po Periodic Table - Extended Periodic Table
General Name, Symbol, Number
tellurium, Te, 52
Element category
metalloids
Group, Period, Block
16, 5, p
Appearance
silvery lustrous gray
Standard atomic weight
127.60(3) g·mol
Electron configuration
[Kr] 5s 4d
−1
2
Electrons per shell
10
4
5p
2, 8, 18, 18, 6 Physical properties
Phase
solid
Density (near r.t.)
6.24 g·cm−3
Liquid density at m.p.
5.70 g·cm−3
Melting point
722.66 K (449.51 °C, 841.12 °F)
Boiling point
1261 K (988 °C, 1810 °F)
Heat of fusion
17.49 kJ·mol−1
Heat of vaporization
114.1 kJ·mol−1
Specific heat capacity
(25 °C) 25.73 J·mol−1·K−1 Vapor pressure
P(Pa)
1
at T(K)
10
100
1k
10 k
100 k
(775)
(888)
1042
1266
Atomic properties Crystal structure
hexagonal
Tellurium
2
Oxidation states
6, 5, 4, 2, -2 (mildly acidic oxide)
Electronegativity
2.1 (Pauling scale)
Ionization energies (more)
1st: 869.3 kJ·mol−1 2nd: 1790 kJ·mol−1 3rd: 2698 kJ·mol−1
Atomic radius
140 pm
Covalent radius
138±4 pm
Van der Waals radius
206 pm Miscellaneous [1]
Magnetic ordering
diamagnetic
Thermal conductivity
(300 K) −1 −1 (1.97–3.38) W·m ·K
Speed of sound (thin rod)
(20 °C) 2610 m/s
Young's modulus
43 GPa
Shear modulus
16 GPa
Bulk modulus
65 GPa
Mohs hardness
2.25
Brinell hardness
180 MPa
CAS registry number
13494-80-9 Most-stable isotopes
Main article: Isotopes of tellurium iso 120
Te
121
Te
NA 0.09% syn
half-life >2.2×1016y 16.78 d
122
2.55%
123
0.89%
124
4.74%
124
125
7.07%
125
126
18.84%
126
Te Te Te Te Te
127
Te
128
Te
129
Te
130
Te
syn 31.74% syn 34.08%
DM
DE (MeV)
DP
εε
1.701
120
ε
1.040
121
0.051
123
Sn Sb
122
Te is stable with 70 neutron
>1.0×1013 y
ε
Sb
Te is stable with 72 neutron Te is stable with 73 neutron Te is stable with 74 neutron
9.35 h
β−
0.698
127
2.2×1024 y
β−β−
0.867
128
69.6 min
β−
1.498
129
7.9×1020 y
β−β−
2.528
130
References
I Xe I Xe
Tellurium
3
Tellurium (pronounced /tɪˈlʊəriəm, tɛlˈlʊəriəm/) is a chemical element that has the symbol Te and atomic number 52. A brittle silver-white metalloid which looks similar to tin, tellurium is chemically related to selenium and sulfur. Tellurium is primarily used in alloys and as a semiconductor.
Characteristics Tellurium is extremely rare, one of the nine rarest metallic elements on Earth. It is in the same chemical family as oxygen, sulfur, selenium, and polonium (the chalcogens). When crystalline, tellurium is silvery-white and when it is in pure state it has a metallic luster. This is a brittle and easily pulverized metalloid. Amorphous tellurium is found by precipitating it from a solution of tellurous or telluric acid (Te(OH)6). However, there is some debate whether this form is really amorphous or made of minute crystals.
Applications Tellurium is a p-type semiconductor that shows a greater conductivity in certain directions which depends on atomic alignment. Chemically related to selenium and sulfur, the conductivity of this element increases slightly when exposed to light (photoconductivity). It can be doped with copper, gold, silver, tin, or other metals. When in its molten state, tellurium is corrosive to copper, iron, and stainless steel. Tellurium gives a greenish-blue flame when burned in normal air and forms tellurium dioxide as a result. Metal alloys:
[2]
• It is mostly used in alloys with other metals. It is added to lead to improve its strength and durability, and to decrease the corrosive action of sulfuric acid. • When added to stainless steel and copper it makes these metals more workable. It is alloyed into cast iron for chill control. Other uses: • Used in ceramics. • It is used in chalcogenide glasses. • Tellurium is used in blasting caps. • Organic tellurides have been employed as initiators for living radical polymerisation and electron-rich mono- and di-tellurides possess antioxidant activity. • Tellurite agar is used to identify member of the corynebacterium genus, most typically Corynebacterium diphtheriae, the pathogen responsible for diphtheria. High purity metalorganics of both selenium and tellurium are used in the semiconductor industry, and are prepared by adduct purification.[3] [4] Semiconductor and electronic industry uses: • Tellurium as a tellurium suboxide is used in the media layer of several types of rewritable optical discs, including ReWritable Compact Discs (CD-RW), ReWritable Digital Video Discs (DVD-RW) and ReWritable Blu-ray Discs. [5] [6] • Tellurium is used in the new phase change memory chips[7] developed by Intel.[8] • Bismuth telluride (Bi2Te3) is used in thermoelectric devices.
Tellurium
4
• Tellurium is used in cadmium telluride (CdTe) solar panels. NREL lab tests using this material achieved some of the highest efficiencies for solar cell electric power generation. First Solar started massive commercial production of CdTe solar panels in recent years, significantly increased tellurium demand. If some of the cadmium in CdTe is replaced by zinc then CdZnTe is formed which is used in solid-state x-ray detectors. • Alloyed with both cadmium and mercury, to form mercury cadmium telluride, an infrared sensitive semiconductor material is formed. Organotellurium compounds such as dimethyl telluride, diethyl telluride, diisopropyl telluride, diallyl telluride and methyl allyl telluride are used as precursors for MOVPE growth of II-VI compound semiconductors. Diisopropyl telluride (DIPTe) is employed as the preferred precursor for achieving the low temperature growth of CdHgTe by MOVPE.
History Tellurium (Latin tellus meaning "earth") was discovered in 1782 by the Hungarian Franz-Joseph Müller von Reichenstein (Müller Ferenc) in Nagyszeben (now, Sibiu) Transylvania. In 1789, another Hungarian scientist, Pál Kitaibel, also discovered the element independently, but later he gave the credit to Müller. In 1798, it was named by Martin Heinrich Klaproth who earlier isolated it.[9] Tellurium was used as a chemical bonder in the making of the outer shell of the first atom bomb. The 1960s brought growth in thermoelectric applications for tellurium, as well as its use in free-machining steel, which became the dominant use.
Occurrence With an abundance in the Earth's crust even lower than platinum, tellurium is, apart from the precious metals, the rarest stable solid element in the Earth's crust. Its abundance in the Earth's crust is 1 to 5 ppb, compared with 5 to 37 ppb for platinum. By comparison, even the rarest of the lanthanides have crustal abundances of 500 ppb. The extreme rarity of tellurium in the Earth's crust is not a reflection of its cosmic abundance, which is in fact greater than that of rubidium[10], even though rubidium is ten thousand times more abundant in the Earth's crust. The extraordinarily low abundance of tellurium on Earth is because during the Earth's formation, the stable form of elements in the absence of oxygen and water was controlled by the oxidation and reduction of hydrogen. Under this scenario elements such as tellurium which form volatile hydrides were severely depleted during the formation of the Earth's crust through evaporation. Tellurium and selenium are the heavy elements mostly depleted in the Earth's crust by this process. Tellurium on quartz (Moctezuma, Sonora, Mexico)
Tellurium is sometimes found in its native (elemental) form, but is more often found as the tellurides of gold (calaverite, krennerite, petzite, sylvanite, and others). Tellurium compounds are the most common chemical compounds of gold found in nature (rare non-tellurides such as gold aurostibite and bismuthide are known). Tellurium is also found combined with elements other than gold, in salts of other metals. The principal source of
Tellurium tellurium is from anode sludges produced during the electrolytic refining of blister copper. It is a component of dusts from blast furnace refining of lead. Treatment of 500 tons of copper ore typically yields [11] one pound of tellurium. Tellurium is produced mainly in the United States, Canada, Peru, and Japan. See here [12]. Commercial-grade tellurium is usually marketed as minus 200-mesh powder but is also available as slabs, ingots, sticks, or lumps. The year-end price for tellurium in 2000 was US$14 per pound. In recent years, tellurium price was driven up [13] by increased demand and limited supply, reaching as high as US$100 per pound in 2006. See also here [14]. See also: Telluride, Colorado, category:Telluride minerals
Compounds Tellurium is in the same series as sulfur and selenium and forms similar compounds. A compound with metal or hydrogen and similar ions is called a telluride. Gold and silver tellurides are considered good ores. Compounds with tellurate ion complexes TeO2−4 or TeO6−6 are known as tellurates. Compounds with the anion TeO2−3 are called tellurites. The tellurium analogues of alcohols and thiols, having the functional group –TeH, are called tellurols. The -TeH functional group is also referred to with the prefix tellanyl-. See also: Category:Tellurium compounds
Isotopes There are 30 known isotopes of tellurium with atomic masses that range from 108 to 137. Naturally found tellurium consists of eight isotopes (listed in the main article); three of them are observed to be radioactive. 128Te has the longest known half-life, 2.2×1024 years[15] , among all radioisotopes.[16] Tellurium is the lightest element known to undergo alpha decay, with isotopes 106Te to 110Te being able to undergo this mode of decay.
Precautions Tellurium and tellurium compounds are considered to be mildly toxic and need to be handled with care, although acute poisoning is rare.[17] Tellurium is not reported to be carcinogenic.[17] Humans exposed to as little as 0.01 mg/m3 or less in air develop "tellurium breath", which has a garlic-like odor.[18] The garlic odor that is associated with human intake of tellurium compounds is caused from the tellurium being metabolized by the body. When the body metabolizes tellurium in any oxidation state, the tellurium gets converted into dimethyl telluride, (CH3)2Te, which is volatile and is the cause of the garlic-like smell. Even though the metabolic pathways of tellurium are not known, it is generally assumed that they resemble those of the more extensively studied selenium, because the final methylated metabolic products of the two elements are similar.
5
Tellurium
6
External links • • • •
WebElements.com – Tellurium [19] USGS Mineral Information on Selenium and Tellurium Selenium Tellurium Development Association [21] Comprehensive Data on Tellurium [22]
[20]
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] George, Micheal W. (2007). " Mineral Yearbook 2007: Selenium and Tellurium (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ selenium/ myb1-2007-selen. pdf)". United States geological Survey. . [3] "Ultra-pure organotellurium precursors for the low temperature MOVPE growth of II/VI compound semiconductors". Journal of Crystal Growth 93: 744–749. 1988. doi: 10.1016/0022-0248(88)90613-6 (http:/ / dx. doi. org/ 10. 1016/ 0022-0248(88)90613-6). [4] U.S. Patent 5117021 (http:/ / www. google. com/ patents?vid=5117021) Method for purification of tellurium and selenium alkyls [5] Farivar, Cyrus (2006-10-19). " Panasonic says that its 100GB Blu-ray discs will last a century (http:/ / www. engadget. com/ 2006/ 10/ 19/ panasonic-says-that-its-100gb-blu-ray-discs-will-last-a-century/ )". . Retrieved 2008-11-13. [6] Kenichi Nishiuchi, Hideki Kitaura, Noboru Yamada and Nobuo Akahira (1998). "Dual-Layer Optical Disk with Te–O–Pd Phase-Change Film". Jpn. J. Appl. Phys. 37: 2163-2167. doi: 10.1143/JJAP.37.2163 (http:/ / dx. doi. org/ 10. 1143/ JJAP. 37. 2163). [7] Hudgens, S.; Johnson, B. (2004). " Overview of Phase-Change Chalcogenide Nonvolatile Memory Technology (http:/ / www. engr. sjsu. edu/ sgleixner/ mate270/ LectureNotes/ Hudgens_MRS. pdf)". Material Research Society Bulletin 29 (11): 1–4. . [8] Geppert, Linda. " The New Indelible Memories (http:/ / www. spectrum. ieee. org/ print/ 1501)". IEEE spectrum online. . Retrieved 2009-02-08. [9] Diemann, Ekkehard; Müller, Achim; Barbu, Horia (2002). "Die spannende Entdeckungsgeschichte des Tellurs (1782 - 1798) Bedeutung und Komplexität von Elemententdeckungen". Chemie in unserer Zeit 36 (5): 334–337. doi: 10.1002/1521-3781(200210)36:5<334::AID-CIUZ334>3.0.CO;2-1 (http:/ / dx. doi. org/ 10. 1002/ 1521-3781(200210)36:5<334::AID-CIUZ334>3. 0. CO;2-1). [10] http:/ / www. orionsarm. com/ science/ Abundance_of_Elements. html [11] http:/ / www. stda. net/ se-te. htm [12] http:/ / www. mmta. co. uk/ economicsFacts/ Articles/ MiningJournalReview/ Tellurium. pdf [13] http:/ / arizonageology. blogspot. com/ 2007/ 05/ arizona-tellurium-rush. html [14] http:/ / www. resourceinvestor. com/ pebble. asp?relid=31031 [15] " WWW Table of Radioactive Isotopes: Tellurium (http:/ / ie. lbl. gov/ toi/ nuclide. asp?iZA=520128)". Nuclear Science Division, Lawrence Berkeley National Laboratory. 2008. . [16] " Noble Gas Research (http:/ / presolar. wustl. edu/ work/ noblegas. html#tellurium)". Laboratory for Space Sciences, Washington University in St. Louis. 2008. . [17] Harrison, W; S Bradberry, J Vale (1998-01-28). " Tellurium (http:/ / www. intox. org/ databank/ documents/ chemical/ tellur/ ukpid84. htm)" (HTML). International Programme on Chemical Safety. . Retrieved 2007-01-12. [18] " Tellurium (http:/ / periodic. lanl. gov/ elements/ 52. html)" (HTML). Los Alamos National Laboratory. 2003-12-15. . Retrieved 2007-01-12. [19] [20] [21] [22]
http:/ / www. webelements. com/ webelements/ elements/ text/ Te/ index. html http:/ / minerals. er. usgs. gov/ minerals/ pubs/ commodity/ selenium http:/ / www. stda. net/ home. htm http:/ / www. mrteverett. com/ Chemistry/ pdictable/ q_elements. asp?Symbol=Te
Article Sources and Contributors
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7
Iodine
1
Iodine 53
tellurium ← iodine → xenon
Br ↑
I ↓
At Periodic Table - Extended Periodic Table
General Name, Symbol, Number
iodine, I, 53
Element category
halogens
Group, Period, Block
17, 5, p
Appearance
solid violet-dark gray, lustrous
violet gas
Standard atomic weight
126.90447(3) g·mol−1
Electron configuration
[Kr] 4d10 5s2 5p5
Electrons per shell
2, 8, 18, 18, 7 Physical properties
Phase
solid
Density (near r.t.)
4.933 g·cm−3
Melting point
386.85 K (113.7 °C, 236.66 °F)
Boiling point
457.4 K (184.3 °C, 363.7 °F)
Triple point
386.65 K, 12.1×103 Pa
Critical point
819 K, 11.7 MPa
Heat of fusion
(I ) 15.52 kJ·mol−1 2
Heat of vaporization
(I ) 41.57 kJ·mol−1 2
Iodine
2
Specific heat capacity
(25 °C) (I2) 54.44 J·mol−1·K−1 Vapor pressure (rhombic)
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
260
282
309
342
381
457
Atomic properties Crystal structure
orthorhombic
Oxidation states
7, 5, 3, 1, -1 (strongly acidic oxide)
Electronegativity
2.66 (Pauling scale)
Ionization energies
1st: 1008.4 kJ/mol 2nd: 1845.9 kJ/mol 3rd: 3180 kJ/mol
Atomic radius
140 pm
Covalent radius
139±3 pm
Van der Waals radius
198 pm Miscellaneous [1]
Magnetic ordering
diamagnetic
Electrical resistivity
(0 °C) 1.3×10 Ω·m
Thermal conductivity
(300 K) 0.449 W·m
Bulk modulus
7.7 GPa
7
−1
CAS registry number
−1
·K
7553-56-2 Most-stable isotopes
Main article: Isotopes of iodine iso 123
I
127
I
NA syn 100%
half-life 13 h
DM ε, γ
DE (MeV)
DP
0.16
123
Te
127
I is stable with 74 neutron
129
syn
15.7×106 y
β−
0.194
129
131
syn
8.02070 d
β−, γ
0.971
131
I I
Xe Xe
References
Iodine (pronounced /ˈaɪ.ədaɪn/, English pronunciation: /ˈaɪ.ədɨn/, or in chemistry English pronunciation: /ˈaɪ.ədiːn/; from Greek: ιώδης iodes "violet"), is a chemical element that has the symbol I and atomic number 53. Naturally-occurring iodine is a single isotope with 74 neutrons. Chemically, iodine is the second least reactive of the halogens, and the second most electropositive halogen; trailing behind astatine in both of these categories. However, the element does not occur in the free state in nature. As with all other halogens (members of
Iodine
3
Group XVII in the periodic table), when freed from its compounds iodine forms diatomic molecules (I2). Iodine and its compounds are primarily used in medicine, photography, and dyes. Although it is rare in the solar system and Earth's crust, the iodides are very soluble in water, and the element is concentrated in seawater. This mechanism helps to explain how the element came to be required in trace amounts by all animals and some plants, being the heaviest element commonly used by living organisms (only tungsten, used in enzymes by a few bacteria, is heavier[2] [3] ).
Characteristics Iodine under standard conditions is a shiny grey solid. It can be seen apparently sublimating at standard temperatures into a violet-pink gas that has an irritating odor. This halogen forms compounds with many elements, but is less reactive than the other members of its Group VII (halogens) and has some metallic light reflectance. Elemental iodine dissolves easily in chloroform and carbon tetrachloride. The solubility of elemental iodine in water can be vastly increased by the addition of potassium iodide. The molecular iodine reacts reversibly with the negative ion, creating the triiodide anion, I3−, which dissolves well in water. This is also the formulation of some types of medicinal (antiseptic) iodine, although tincture of iodine classically dissolves the element in alcohol. The deep blue color of starch-iodine complexes is produced only by the free element. Students who have seen the classroom demonstration in which iodine crystals are gently heated in a test tube to violet vapor may gain the impression that liquid iodine does not exist at atmospheric pressure. This misconception arises because the vapor produced has such a deep colour that the liquid appears not to form. In fact, if iodine crystals are heated carefully to just above their melting point of 113.7 °C, the crystals melt into a liquid which is present under a dense blanket of the vapor. When iodine is encapsulated into carbon nanotubes it forms atomic chains, whose structure depends on the nanotube diameter.[4]
Occurrence Iodine naturally occurs in the environment chiefly as a dissolved iodide in seawater, although it is also found in some minerals and soils.[5] This element also exists in small amounts in the mineral caliche, found in Chile, between the Andes and the sea. A type of seaweed, kelp, tends to be high in iodine as well. Organoiodine compounds are produced by marine life forms, the most notable being iodomethane (commonly called methyl iodide). The total iodomethane that is produced by Iodomethane the marine environment, by microbial activitiy in rice paddies and by the burning of biological material is estimated to be 214 kilotonnes.[6] The volatile iodomethane is broken up by oxidation reactions in the atmosphere and a global iodine cycle is established.[5] [6] Although the element is actually quite rare, kelp and certain plants and other algae have some ability to concentrate iodine, which helps introduce the element into the food chain.
Iodine
4
Structure Iodine crystallizes in the orthorombic space group Cmca No 64, Pearson symbol oS8, the same as black phosphorus. In the solid state, I2 molecules are still represented by a short I-I bond of 270 pm.
Production From the several places in which iodine occures in nature only two are used as source for iodine: the caliche, found in Chile and the iodine containing brines of gas and oil fields, especially in Japan and the United States.
Structure of solid iodine
The caliche, found in Chile contains sodium nitrate, which is the main product of the mining activities and small amounts of sodium iodate and sodium iodide. During leaching and production of pure sodium nitrate the sodium iodate and iodide is extracted.[7] The high concentration of iodine in the caliche and the extensive mining made Chile the largest producer of iodine in 2007. Most other producers use natural occurring brine for the production of iodine. The Japanese Minami Kanto gas field east of Tokyo and the American Anadarko Basin gas field in northwest Oklahoma are the two largest sources for iodine from brine. The brine has a temperature of over 60°C due to the depth of the Iodine output in 2005 source. The brine is first purified and acidified using sulfuric acid, then the iodide present is oxidized to iodine with chlorine. An iodine solution is produced, but is dilute and must be concentrated. Air is blown into the solution, causing the iodine to evaporate, then it is passed into an absorbing tower containing acid where sulfur dioxide is added to reduce the iodine. The hydrogen iodide (HI) is reacted with chlorine to precipitate the iodine. After filtering and purification the iodine is packed.[7] [8] 2 HI + Cl2 → I2↑ + 2 HCl I2 + 2 H2O + SO2 → 2 HI + H2SO4 2 HI + Cl2 → I2↓ + 2 HCl The production of iodine from seawater via electrolysis is not used due to the sufficient abundance of iodine-rich brine. Another source of iodine was kelp, used in the 18th and 19th centuries, but no longer economically viable. Commercial samples often contain a large amount of impurities; they may be removed by sublimation. The element may also be prepared in an ultrapure form through the reaction of potassium iodide with copper(II) sulfate, which gives copper(II) iodide initially. That decomposes spontaneously to copper(I) iodide and iodine: Cu2+ + 2 I− → CuI2 2 CuI2 → 2 CuI + I2 There are also other methods of isolating this element in the laboratory, for example the method used to isolate other halogens: oxidation of the iodide in hydroiodic acid (often made in situ with an iodide and sulfuric acid) by manganese dioxide (see below in
Iodine
5
Descriptive chemistry).
Isotopes There are 37 known (characterized) isotopes of iodine, but only one,
127
I, is stable.
In many ways, 129I is similar to 36Cl. It is a soluble halogen, fairly non-reactive, exists mainly as a non-sorbing anion, and is produced by cosmogenic, thermonuclear, and in-situ reactions. In hydrologic studies, 129I concentrations are usually reported as the ratio of 129I to total I (which is virtually all 127I). As is the case with 36Cl/Cl, 129I/I ratios in nature are quite small, 10−14 to 10−10 (peak thermonuclear 129I/I during the 1960s and 1970s reached about 10−7). 129I differs from 36Cl in that its halflife is longer (15.7 vs. 0.301 million years), it is highly biophilic, and occurs in multiple ionic forms (commonly, I− and IO3−) which have different chemical behaviors. This makes it fairly easy for 129I to enter the biosphere as it becomes incorporated into vegetation, soil, milk, animal tissue, etc. Excesses of stable 129Xe in meteorites have been shown to result from decay of "primordial" iodine-129 produced newly by the supernovas which created the dust and gas from which the solar system formed. 129I was the first extinct radionuclide to be identified as present in the early solar system. Its decay is the basis of the I-Xe Iodine-xenon radiometric dating scheme, which covers the first 85 million years of solar system evolution. Effects of various radioiodine isotopes in biology are discussed below.
History Iodine was discovered by Bernard Courtois in 1811.[9] [10] He was born to a manufacturer of saltpeter (a vital part of gunpowder). At the time of the Napoleonic Wars, France was at war and saltpeter was in great demand. Saltpeter produced from French niter beds required sodium carbonate, which could be isolated from seaweed washed up on the coasts of Normandy and Brittany. To isolate the sodium carbonate, seaweed was burned and the ash then washed with water. The remaining waste was destroyed by adding sulfuric acid. One day Courtois added too much sulfuric acid and a cloud of purple vapor rose. Courtois noted that the vapor crystallized on cold surfaces making dark crystals. Courtois suspected that this was a new element but lacked the money to pursue his observations. However he gave samples to his friends, Charles Bernard Desormes (1777–1862) and Nicolas Clément (1779–1841), to continue research. He also gave some of the substance to Joseph Louis Gay-Lussac (1778–1850), a well-known chemist at that time, and to physicist André-Marie Ampère (1775–1836). On 29 November 1813, Dersormes and Clément made public Courtois’s discovery. They described the substance to a meeting of the Imperial Institute of France. On December 6, Gay-Lussac announced that the new substance was either an element or a compound of oxygen.[11] [12] [13] Ampère had given some of his sample to Humphry Davy (1778–1829). Davy did some experiments on the substance and noted its similarity to chlorine.[14] Davy sent a letter dated December 10 to the Royal Society of London stating that he had identified a new element.[15] A large argument erupted between Davy and Gay-Lussac over who identified iodine first but both scientists acknowledged Courtois as the first to isolate the chemical element.
Iodine
Applications Disinfectant and water treatment Elemental iodine is used as a disinfectant in various forms. The iodine exists as the element, or as the water soluble triiodide anion I3- generated in situ by adding iodide to poorly-soluble iodine (the reverse chemical reaction makes some free elemental iodine availalbe for antisepsis). Alternatively, iodine may come from iodophors, which contain iodine complexed with a solubilizing agent (iodide ion may be thought of loosely as the iodophor in triiodide water solutions). Examples of such preparations include:[16] • Tincture of iodine (iodine in ethanol, or iodine and sodium iodide in a mixture of ethanol and water) • Lugol's iodine (iodine and iodide in water, forming mostly tiiodide) • Povidone iodine (an iodophor)
Staining Iodine is a common general stain used in thin-layer chromatography. It is also used in the Gram stain as a mordant, after the sample is treated with crystal violet. In particular, iodine forms an intense blue complex with the glucose polymers starch andglycogen. Many applications rely on this property: • Iodometry. The concentration of an oxidant can be determined by adding it to an excess of iodide with a little free iodine, to destroy elemental iodine/triiodide Testing a seed for starch with a as a result of oxidation by the oxidant. A starch solution of iodine indicator is then used as the indicator close to the end-point, in order to increase the visual contrast (dark blue becomes colorless, instead of the yellow of dilute triiodide becoming colorless). • An Iodine test may be used to test a sample substance for the presence of starch. • The Iodine clock reaction is an extension of the techniques in iodometry. • Iodine solutions are used in counterfeit banknote detection pens; the premise being that counterfeit banknotes made using commercially available paper contain starch. • Starch-iodide paper are used to test for the presence of oxidants such as peroxides. The oxidants convert iodide to iodine, which shows up as blue. A solution of starch and iodide can perform the same function.[17] • During colposcopy, Lugol's iodine is applied to the vagina and cervix. Normal vaginal tissue stains brown due to its high glycogen content (a color-reaction similar to that with starch), while abnormal tissue suspicious for cancer does not stain, and thus appears pale compared to the surrounding tissue. Biopsy of suspicious tissue can then be performed. This is called a Schiller's test.
6
Iodine
7
Radiocontrast agent Iodine, as a heavy element, is quite radio-opaque. Organic compounds of a certain type (typically iodine-substituted benzene derivatives) are thus used in medicine as X-ray radiocontrast agents for intravenous injection. This is often in conjunction with advanced X-ray techniques such as angiography and CT scanning
Radioiodine Some radioactive iodine isotopes can be used to treat Diatrizoic acid, a radiocontrast thyroid cancer. The body accumulates iodine in the thyroid, thus radioactive iodine can selectively damage growing thyroid cancer cells while the radioactive dose to the rest of the body remains small.
Iodine compounds Iodine forms many compounds. Potassium iodide is the most commercially significant iodine compound. It is a convenient source of the iodide anion; it is easier to handle than sodium iodide because it is not hygroscopic. Sodium iodide is especially useful in the Finkelstein reaction, because it is soluble in acetone, while potassium iodide is poorly so. In this reaction, an alkyl chloride is converted to an alkyl iodide. This relies on the insolubility of sodium chloride in acetone to drive the reaction: R-Cl
(acetone)
+ NaI
(acetone)
→ R-I
(acetone)
+ NaCl
(s)
Iodic acid (HIO3) and its salts are strong oxidizers. Periodic acid (HIO4) cleaves vicinal diols along the C-C bond to give aldehyde fragments. 2-Iodoxybenzoic acid and Dess-Martin periodinane are hypervalent iodine oxidants used to specifically oxidize alcohols to ketones or aldehydes. Iodine pentoxide is a strong oxidant as well. Interhalogen compounds are well known; examples include iodine monochloride and trichloride; iodine pentafluoride and heptafluoride. HI
He BeI2
BI3
CI4
NI3
I2O4, I2O5, I4O9
IF, IF3, IF5, IF7
Ne
NaI MgI2
AlI3
SiI4
PI3, P2I4
S
ICl, ICl3
Ar
LiI
KI
CaI2
Sc
TiI4
VI3
RbI
SrI2
Y
ZrI4
CsI
BaI2
Fr
Ra
Cr
CuI
ZnI2
Ga2I6
GeI2, GeI4
AsI3
Se
IBr
Kr
Pd
AgI
CdI2
InI3
SnI4, SnI2
SbI3
TeI4
I
Xe
Ir
Pt
AuI
Hg2I2, HgI2
TlI
PbI2
Bi
Po
At
Rn
Mt
Ds
Rg
Uub
Uut
Uuq
Uup
Uuh
Uus
Uuo
MnI2
Fe
CoI2 NiI2
Nb Mo
Tc
Ru
Rh
Hf
Ta
W
Re
Os
Rf
Db
Sg
Bh
Hs
Iodine
8 ↓ La
Ce
Ac ThI4
Pr
Nd
Pm
SmI2
Eu
Gd
TbI3
Dy
Ho
Er
Tm
Yb
Lu
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
Organic compounds Many organoiodine compounds exist, the simplest is iodomethane, approved as a soil fumigant. Iodinated organics are used as synthetic reagents, and also radiocontrast agents. Biologically active substances organoiodine compounds.[18]
like
the
thyroid
hormones
are
naturally
occurring
Chemistry Elemental iodine is poorly soluble in water, with one gram dissolving in 3450 ml at 20 °C and 1280 ml at 50 °C. By contrast with chlorine, the formation of the hypohalite ion (IO–) in neutral aqueous solutions of iodine is negligible. I2+ H2O
H+ + I− + HIO (K = 2.0×10−13)[19]
Solubility in water is greatly improved if the solution contains dissolved iodides such as hydroiodic acid, potassium iodide, or sodium iodide; this extra solubility results from the high solubility of the I3− ion. Dissolved bromides also improve water solubility of iodine. Iodine is soluble in a number of organic solvents, including ethanol (20.5 g/100 ml at 15 °C, 21.43 g/100 ml at 25 °C), diethyl ether (20.6 g/100 ml at 17 °C, 25.20 g/100 ml at 25 °C), chloroform, acetic acid, glycerol, benzene (14.09 g/100 ml at 25 °C), carbon tetrachloride (2.603 g/100 ml at 35 °C), and carbon disulfide (16.47 g/100 ml at 25 °C).[20] Aqueous and ethanol solutions are brown. Solutions in chloroform, carbon tetrachloride, and carbon disulfide are violet. Molecular iodine can be prepared by oxidizing iodides with chlorine: 2 I− + Cl2 → I2 + 2 Cl− or with manganese dioxide in acid solution:[19] 2 I− + 4 H+ + MnO2 → I2 + 2 H2O + Mn2+ Iodine is reduced to hydroiodic acid by hydrogen sulfide:[21] I2 + H2S → 2 HI + S↓ or by hydrazine: 2 I2 + N2H4 → 4 HI + N2 Iodine is oxidized to iodate by nitric acid:[22] I2 + 10 HNO3 → 2 HIO3 + 10 NO2 + 4 H2O or by chlorates:[22] I2 + 2 ClO3− → 2 IO3− + Cl2 Iodine is converted in a two stage reaction to iodide and iodate in solutions of alkali hydroxides (such as sodium hydroxide):[19] I2 + 2 OH− → I− + IO− + H2O
(K = 30)
3 IO− → 2 I− + IO3−
(K = 1020)
Iodine
9
Despite having the lowest electronegativity of the common halogens, iodine reacts violently with some metals, such as aluminum: 3 I2 + 2 Al → 2 AlI3 This reaction produces 314 kJ per mole of aluminum, comparable to thermite's 425 kJ. Yet the reaction initiates spontaneously, and if unconfined, causes a cloud of gaseous iodine due to the high heat.
Organic synthesis With phosphorus, iodine is able to replace hydroxyl groups on alcohols with iodide. For example, the synthesis of methyl iodide from methanol, red phosphorus, and iodine.[23] The iodinating reagent is phosphorus triiodide that is formed in situ: 3 CH3OH + PI3 → 3 CH3I + H3PO3 Phosphorous acid is formed as a side-product. The iodoform test uses an alkaline solution of iodine to react with methyl ketones to give the labile triiodomethide leaving group, forming iodoform which precipitates. Iodine is sometimes used to activate magnesium when preparing Grignard reagents; aryl and alkyl iodides both form Grignard reagents. Alkyl iodides such as iodomethane are good alkylating agents. Some drawbacks to use of iodo-organics in chemical synthesis are: • iodine compounds tend to be more expensive than the corresponding bromides and chlorides, in that order • iodides tend to be much stronger alkylating agents, and so are more toxic (e.g. methyl iodide is very toxic (T+)[24] • low molecular weight iodides tend to have a much higher equivalent weight, compared with other alkylating agents (e.g. methyl iodide versus dimethyl carbonate), due to the atomic mass of iodine.
Clandestine synthetic chemical use In the United States, the Drug Enforcement Agency (DEA) regards iodine and compounds containing iodine (ionic iodides, iodoform, ethyl iodide, and so on) as reagents useful for the clandestine manufacture of methamphetamine. Persons who attempt to purchase significant quantities of such chemicals without establishing a legitimate use are likely to find themselves the target of a DEA investigation. Persons selling such compounds without doing due diligence to establish that the materials are not being diverted to clandestine use may be subject to stiff penalties, such as expensive fines or even imprisonment.[25] [26]
Biological role Iodine is an essential trace element for life, the heaviest element commonly needed by living organisms, and the second-heaviest known to be used by any form of life (only tungsten, a component of a few bacterial enzymes, has a higher atomic number and atomic weight). Iodine's main role in animal biology is as constituents of the thyroid hormones, thyroxine (T4) and triiodothyronine (T3). These are made from addition condensation products of the amino acid tyrosine, and are stored prior to release in an iodine-containing protein called thyroglobulin. T4 and T3 contain four and three atoms of iodine per molecule, respectively. The thyroid gland actively absorbs iodide from the blood to make and release these hormones into the blood, actions which are regulated by a second hormone TSH from
Iodine
10
the pituitary. Thyroid hormones are phylogenetically very old molecules which are synthesized by most multicellular organisms, and which even have some effect on unicellular organisms. Thyroid hormones play a basic role in biology, acting on gene transcription to regulate the basal metabolic rate. The total deficiency of thyroid hormones can reduce basal metabolic rate up to 50%, while in excessive production of thyroid hormones the basal metabolic rate can be increased by 100%. T4 acts largely as a precursor to T3, which is (with minor exceptions) the biologically active hormone.
Extrathyroidal Iodine Iodine accounts for 65% of the molecular weight of T4 and 59% of the T3. 15–20 mg of iodine is concentrated in thyroid tissue and hormones, but 70% of the body's iodine is distributed in other tissues, including mammary glands, eyes, gastric mucosa, the cervix, and salivary glands. In the cells of these tissues iodide enters direcly by sodium-iodide symporter (NIS). Its role in mammary tissue is related to fetal and neonatal development, but its role in the other tissues is unknown.[27] It has been shown to act as an antioxidant in these tissues.[27] Iodine may have a relationship with selenium, and iodine supplementation selenium-deficient populations may pose risks for thyroid function.[27]
in
The US Food and Nutrition Board and Institute of Medicine recommended daily allowance of iodine ranges from 150 micrograms /day for adult humans to 290 micrograms /day for lactating mothers. However, the thyroid gland needs no more than 70 micrograms /day to synthesize the requisite daily amounts of T4 and T3. These higher recommended daily allowance levels of iodine seem necessary for optimal function of a number of body systems, including lactating breast, gastric mucosa, salivary glands, oral mucosa, thymus, epidermis, choroid plexus, etc.[28] [29] [30] [31]
Iodine and cancer *Breast cancer. It is known that a diet lacking in iodine is connected with adverse health effects collectively referred as iodine deficiency diseases or disorders. Studies also indicate that iodine deficiency, either dietary or pharmacologic, can lead to breast atypia and increased incidence of malignancy in animal models, while iodine treatment can reverse dysplasia.[32] [33] [34] Laboratory evidences demonstrate that the effect of iodine on breast cancer is in part independent of thyroid function and that iodine inhibit cancer promotion through modulation of the estrogen pathway. Gene array profiling of estrogen responsive breast cancer cell line shows that the combination of iodine and iodide alters gene expression and inhibits the estrogen response through up-regulating proteins involved in estrogen metabolism. Whether iodine/iodide will be useful as an adjuvant therapy in the pharmacologic manipulation of the estrogen pathway in women with breast cancer has not been determined clinically.[32] *Iodine and stomach cancer Some researchers have found a epidemiologic correlation between iodine deficiency, iodine-deficient goitre and gastric cancer;[35] [36] [37] a decrease of the incidence of death rate from stomach cancer after implementation of the effective iodine-prophylaxis was reported too.[38] The proposed mechanism of action is that iodide ion can function in gastric mucosa as an antioxidant reducing species that can detoxify poisonous reactive oxygen species, such as hydrogen peroxide.
Iodine
Iodine and immunity Iodine has important actions in the immune system. The high iodide-concentration of thymus suggests an anatomical rationale for this role of iodine in immune system. [39] [40] [41] [42] [43] [44]
Human dietary intake The United States Recommended Daily Allowance (RDA) is 150 micrograms per day (μg/day) for both men and women, with a Tolerable Upper Intake Level (UL) for adults is 1,100 μg/day (1.1 mg/day).[45] The tolerable upper limit was assessed by analyzing the effect of supplementation on thyroid-stimulating hormone.[27] Natural sources of iodine include sea life, such as kelp and certain seafood, as well as plants grown on iodine-rich soil.[46] [47] Iodized salt is fortified with iodine.[47] As of 2000, the median intake of iodine from food in the United States was 240 to 300 μg/day for men and 190 to 210 μg/day for women.[45] In Japan, consumption is much higher due to the frequent consumption of seaweed or kombu kelp.[27] After iodine fortification programs (e.g. iodized salt) have been implemented, some cases of iodine-induced hyperthyroidism have been observed (so called Jod-Basedow disease). The condition mainly seems to occur in people over forty, and the risk appears higher when iodine deficiency is severe and the initial rise in iodine intake is high.[48]
Deficiency In areas where there is little iodine in the diet, typically remote inland areas and semi-arid equatorial climates where no marine foods are eaten, iodine deficiency gives rise to hypothyroidism, symptoms of which are extreme fatigue, goitre, mental slowing, depression, weight gain, and low basal body temperatures.[49] Iodine deficiency is the leading cause of preventable mental retardation, a result which occurs primarily when babies or small children are rendered hypothyroidic by a lack of the element. The addition of iodine to table salt has largely eliminated this problem in the wealthier nations, but as of March 2006, iodine deficiency remained a serious public health problem in the developing world.[50] Iodine deficiency is also a problem in certain areas of Europe. In Germany it has been estimated to cause a billion dollars in healthcare costs per year.[27]
Radioiodine in biology Radioiodine and the thyroid The most common compounds of iodine are the iodides of sodium (NaI) and potassium (KI) and the iodates (KIO3), as elemental iodine is mildly toxic to all living things. Normal iodine is an essential precursor for the manufacture of thyroid hormone. Due to preferential uptake of iodine by the thyroid, isotopes with short half lives such was I131 can be used for thyroid ablation, a procedure in which radioactive iodine is administed intravenously following a diagnostic scan. This procedure is generally performed on patients with thyroid cancer or hyperfunctioning thyroid tissue. After uptake, the iodine undergoes degeneration via beta decay, destroying its associated thyroid tissue. Normally thyroidectomy is preformed prior to ablation to avoid side effects of epilation and radiation
11
Iodine toxicity. The purpose of radioablation is to destroy remnant tissue that was unable to be removed with surgery. Lower energy isotopes such as iodine-123, and less commonly iodine-125, are used as tracers to evaluate the anatomic and physiologic function of the thyroid. Abnormal results may be caused by disorders such as Graves' Disease or Hashimoto's thyroiditis. Potassium iodide tablets have been distributed to populations exposed to nuclear fission accidents such as the Chernobyl disaster. Alternatively, SSKI, a saturated solution of potassium (K) iodide in water, in the form of drops, has been used. In theory, many harmful late-cancer effects of nuclear fallout might be prevented in this way, since an excess of thyroid cancers, presumably due to radioiodine uptake, is the only proven radioisotope contamination effect after a fission accident, or from contamination by fallout from at atomic bomb (prompt radiation from the bomb also cases other cancers, such as leukemias, directly). Taking large amounts of thyroid saturates iodide receptors prevents uptake of most radioactive iodine-131 that may be present from fission product exposure (although it does not protect from other radioisotopes, nor from any other form of direct radiation). The protective effect of KI lasts approximately 24 hours, so must be dosed daily until a risk of significant exposure to radioiodines from fission products no longer exists.[51] [52] Iodine-131 (the most common radioiodine contaminant in fallout) also decays relatively rapidly with a half-life of 8 days, so that 99.95% of the original radioiodine is gone after three months.
Iodine 129 Iodine-129 (129I; half-life 15.7 million years) is a product of cosmic ray spallation on various isotopes of xenon in the atmosphere, in cosmic ray muon interaction with tellurium-130, and also uranium and plutonium fission, both in subsurface rocks and nuclear reactors. Artificial nuclear processes, in particular nuclear fuel reprocessing and atmospheric nuclear weapons tests, have now swamped the natural signal for this isotope. Nevertheless, it now serves as a groundwater tracer as indicator of nuclear waste dispersion into the natural environment. In a similar fasion, 129I was used in rainwater studies to track fission products following the Chernobyl disaster. Radioiodine and the kidney In the 1970s imaging techniques were developed in California to utilize radioiodine in diagnostics for renal hypertension, although this is clinically not commonly performed today and has been placaded by other chemical compounds such as DMSA.
Precautions and toxicity of elemental iodine Elemental iodine is an oxidizing irritant and direct contact with skin can cause lesions, so iodine crystals should be handled with care. Solutions with high elemental iodine concentration such as tincture of iodine are capable of causing tissue damage if use for cleaning and antisepsis is prolonged. Elemental iodine (I2) is poisonous if taken orally in larger amounts; 2–3 grams of it are a lethal dose for an adult human. Iodine vapor is very irritating to the eye, to mucous membranes, and in the respiratory tract. Concentration of iodine in the air should not exceed 1 mg/m³ (eight-hour time-weighted average).
12
Iodine
13
When mixed with ammonia and water, elemental iodine forms nitrogen triiodide which is extremely shock sensitive and can explode unexpectedly.
Toxicity of iodide ion Excess iodine has symptoms similar to those of iodine deficiency. Commonly encountered symptoms are abnormal growth of the thyroid gland and disorders in functioning and growth of the organism as a whole. Iodides are similar in toxicity to bromides.
See also • • • •
Iodide as an antioxidant Chemical Oxygen Iodine Laser Nutrition facts label Starch indicator
External links • "Micronutrient Research for Optimum Health", Linus Pauling Institute, OSU Oregon State University [53] • ATSDR - CSEM: Radiation Exposure from Iodine 131 [54] U.S. Department of Health and Human Services (public domain) • ChemicalElements.com - Iodine [55] • who.int, WHO Global Database on Iodine Deficiency [56] • Network for Sustained Elimination of Iodine Deficiency [57] • Oxidizing Agents > Iodine [58] • WebElements.com – Iodine [59] pnb:نیڈویئآ
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] J McMaster and John H Enemark (1998). "The active sites of molybdenum- and tungsten-containing enzymes". Current Opinion in Chemical Biology 2: 201. doi: 10.1016/S1367-5931(98)80061-6 (http:/ / dx. doi. org/ 10. 1016/ S1367-5931(98)80061-6). [3] Russ Hille (2002). "Molybdenum and tungsten in biology". Trends in Biochemical Sciences 27: 360. doi: 10.1016/S0968-0004(02)02107-2 (http:/ / dx. doi. org/ 10. 1016/ S0968-0004(02)02107-2). [4] Guan, L; Suenaga, K; Shi, Z; Gu, Z; Iijima, S (Jun 2007). "Polymorphic structures of iodine and their phase transition in confined nanospace.". Nano letters 7: 1532. doi: 10.1021/nl070313t (http:/ / dx. doi. org/ 10. 1021/ nl070313t). PMID 17477579. [5] Dissanayake, C. B.; Chandrajith, Rohana; Tobschall, H. J. (1999). "The iodine cycle in the tropical environment — implications on iodine deficiency disorders". International Journal of Environmental Studies 56: 357. doi: 10.1080/00207239908711210 (http:/ / dx. doi. org/ 10. 1080/ 00207239908711210). [6] N. Bell, L. Hsu, D. J. Jacob, M. G. Schultz, D. R. Blake, J. H. Butler, D. B. King, J. M. Lobert, and E. Maier-Reimer (2002). "Methyl iodide: Atmospheric budget and use as a tracer of marine convection in global models". Journal of GeophysicalResearch 107: 4340. doi: 10.1029/2001JD001151 (http:/ / dx. doi. org/ 10. 1029/ 2001JD001151). [7] Jessica Elzea Kogel, Nikhil C. Trivedi, James M. Barker, Stanley T. Krukowski (2006). Industrial Minerals & Rocks: Commodities, Markets, and Uses (http:/ / www. google. com/ books?id=zNicdkuulE4C). SME. pp. 541–552. ISBN 9780873352338. . [8] Tatsuo Maekawa, Shun-Ichiro Igari and Nobuyuki Kaneko (2006). "Chemical and isotopic compositions of brines from dissolved-in-water type natural gas fields in Chiba, Japan". Geochemical Journal 40: 475. doi: 10.2343/geochemj.40.475 (http:/ / dx. doi. org/ 10. 2343/ geochemj. 40. 475).
Iodine [9] Bernard Courtois (1813). "Découverte d'une substance nouvelle dans le Vareck". Annales de chimie 88: 304. In French, seaweed that had been washed onto the shore was called "varec", "varech", or "vareck", whence the English word "wrack". Later, "varec" also referred to the ashes of such seaweed: the ashes were used as a source of iodine and salts of sodium and potassium. [10] Patricia A. Swain (2005). " Bernard Courtois (1777-1838) famed for discovering iodine (1811), and his life in Paris from 1798 (http:/ / www. scs. uiuc. edu/ ~mainzv/ HIST/ awards/ OPA Papers/ 2007-Swain. pdf)". Bulletin for the History of Chemistry 30 (2): 103. . [11] J. Gay-Lussac (1813). "Sur un nouvel acide formé avec la substance décourverte par M. Courtois". Annales de chimie 88: 311. [12] J. Gay-Lussac (1813). "Sur la combination de l'iode avec d'oxigène". Annales de chimie 88: 319. [13] J. Gay-Lussac (1814). "Mémoire sur l'iode". Annales de chimie 91: 5. [14] H. Davy (1813). "Sur la nouvelle substance découverte par M. Courtois, dans le sel de Vareck". Annales de chemie 88: 322. [15] Humphry Davy (January 1, 1814). " Some Experiments and Observations on a New Substance Which Becomes a Violet Coloured Gas by Heat (http:/ / rstl. royalsocietypublishing. org/ content/ 104/ 74. full. pdf+ html)". Phil. Trans. R. Soc. Lond. 104: 74. doi: 10.1098/rstl.1814.0007 (http:/ / dx. doi. org/ 10. 1098/ rstl. 1814. 0007). . [16] Block, Seymour Stanton (2001). Disinfection, sterilization, and preservation. Hagerstwon, MD: Lippincott Williams & Wilkins. p. 159. ISBN 0-683-30740-1. [17] R. Toreki. " Peroxide (http:/ / www. ilpi. com/ msdS/ ref/ peroxide. html)". The MSDS HyperGlossary. . [18] Gribble, G. W. (1996). "Naturally occurring organohalogen compounds - A comprehensive survey". Progress in the Chemistry of Organic Natural Products 68: 1–423. PMID 8795309. [19] F. A. Cotton and G. Wilkinson (1988). Advanced Inorganic Chemistry, 5th ed.. John Wiley & Sons. ISBN 0471849979. [20] Martha Windholz, editor ; Susan Budavari, associate editor ; Lorraine Y. Stroumtsos, assistant editor ; Margaret Noether Fertig, assistant editor. (1976). Merck Index of Chemicals and Drugs, 9th ed. S.l.: J A Majors Company. ISBN 0911910263. [21] N.L. Glinka (1981). General Chemistry (volume 2). Mir Publishing. [22] Linus Pauling (1988). General Chemistry. Dover Publications. ISBN 0486656225. [23] King, C. S.; Hartman, W. W. (1943). " Methyl Iodide (http:/ / www. orgsyn. org/ orgsyn/ orgsyn/ prepContent. asp?prep=CV2P0399)". Org. Synth.; Coll. Vol. 2: 399. [24] " Safety data for iodomethane (http:/ / msds. chem. ox. ac. uk/ IO/ iodomethane. html)". Oxford University. . [25] " 21 USC Sec. 872 01/22/02 (http:/ / www. usdoj. gov/ dea/ pubs/ csa/ 872. htm)". 21. [26] " Chemical Supplier Convicted of Diversion of Iodine (http:/ / www. usdoj. gov/ dea/ pubs/ states/ newsrel/ sanfran032205. html)". . [27] Patrick L (2008). " Iodine: deficiency and therapeutic considerations (http:/ / www. thorne. com/ altmedrev/ . fulltext/ 13/ 2/ 116. pdf)". Altern Med Rev 13: 116. PMID 18590348. . [28] Brown-Grant, K. (1961). " Extrathyroidal iodide concentrating mechanisms (http:/ / physrev. physiology. org/ cgi/ reprint/ 41/ 1/ 189. pdf)". Physiol Rev. 41: 189. . [29] Spitzweg, C., Joba, W., Eisenmenger, W. and Heufelder, A.E. (1998). "Analysis of human sodium iodide symporter gene expression in extrathyroidal tissues and cloning of its complementary deoxyribonucleic acid from salivary gland, mammary gland, and gastric mucosa". J Clin Endocrinol Metab. 83: 1746. doi: 10.1210/jc.83.5.1746 (http:/ / dx. doi. org/ 10. 1210/ jc. 83. 5. 1746). [30] Banerjee, R.K., Bose, A.K., Chakraborty, t.K., de, S.K. and datta, A.G. (1985). "Peroxidase catalysed iodotyrosine formation in dispersed cells of mouse extrathyroidal tissues". J Endocrinol. 2: 159. [31] Miller, D.W. (2006). "Extrathyroidal Benefits of Iodine". J Am Phys and Surg. 11: 106. [32] Stoddard II, F. R.; Brooks, A. D.; Eskin, B. A.; Johannes, G. J. (2008). " Iodine Alters Gene Expression in the MCF7 Breast Cancer Cell Line: Evidence for an Anti-Estrogen Effect of Iodine (http:/ / www. medsci. org/ v05p0189. htm)". International Journal of Medical Science 5: 189. PMID 18645607. . [33] Eskin, B. A.; Grotkowski, C. E.; Connolly, C. P.; Ghent W. R.; (1995). "Different tissue responses for iodine and iodide in rat thyroid and mammary glands". Bioligal Trace Elements Research 49: 9. doi: 10.1007/BF02788999 (http:/ / dx. doi. org/ 10. 1007/ BF02788999). PMID 14965610. [34] Venturi, S. (2001). "Is there a role for iodine in breast diseases?". The Breast 10: 379. doi: 10.1054/brst.2000.0267 (http:/ / dx. doi. org/ 10. 1054/ brst. 2000. 0267). PMID 7577324. [35] Josefssson M, Ekblad E. (2009). Victor R. Preedy, Gerard N. Burrow MD, Ronald Watson. ed. Sodium Iodide Symporter (NIS) in Gastric Mucosa: Gastric Iodide Secretion. In: Comprehensive Handbook of Iodine: Nutritional, Biochemical, Pathological and Therapeutic Aspects. [36] Abnet CC, Fan JH, Kamangar F, Sun XD, Taylor PR, Ren JS, Mark SD, Zhao P, Fraumeni JF Jr, Qiao YL, Dawsey SM (2006). Self-reported goiter is associated with a significantly increased risk of gastric noncardia adenocarcinoma in a large population-based Chinese cohort.. 119. p. 1508.
14
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15
[37] Behrouzian R, Aghdami N. (2004). East Mediterr Health J.. 10. p. 921. [38] Golkowski F, Szybinski Z, Rachtan J, Sokolowski A, Buziak-Bereza M, Trofimiuk M, Hubalewska-Dydejczyk A, Przybylik-Mazurek E, Huszno B. (2007). "Iodine prophylaxis--the protective factor against stomach cancer in iodine deficient areas". Eur J Nutr. 46: 251. doi: 10.1007/s00394-007-0657-8 (http:/ / dx. doi. org/ 10. 1007/ s00394-007-0657-8). [39] Venturi S, Venturi M (September 2009). "Iodine, thymus, and immunity". Nutrition 25 (9): 977–9. doi: 10.1016/j.nut.2009.06.002 (http:/ / dx. doi. org/ 10. 1016/ j. nut. 2009. 06. 002). PMID 19647627. [40] Venturi S.; Venturi A, Cimini D, Arduini C, Venturi M, Guidi A. (1993). "A new hypothesis: iodine and gastric cancer.". Europ. J. Cancer. Prev. 2: 17. [41] Marani L; Venturi S, Masala R (1985). "Role of iodine in delayed immune response.". Isr. J. Med. Sci. 21: 864. [42] Ma F; Zhao W, Kudo M, Aoki K, Misumi J. (2002). "Inhibition of vacuolation toxin activity of Helicobacter pylori by iodine, nitrite and potentiation by sodium chloride, sterigmatocystin and fluoride.". Toxicol in Vitro 16: 531. doi: 10.1016/S0887-2333(02)00045-0 (http:/ / dx. doi. org/ 10. 1016/ S0887-2333(02)00045-0). [43] Klebanoff S.J. (1967). "Iodination of bacteria: A bacterial mechanism.". J Exp Med 126: 1063. doi: 10.1084/jem.126.6.1063 (http:/ / dx. doi. org/ 10. 1084/ jem. 126. 6. 1063). [44] "Iodine enhances ig-G-synthesis by human peripheral blood Iyphocytes in vitro.". Acta Endocr 103: 103. 1983. [45] United States National Research Council (2000). Dietary Reference Intakes for Vitamin A, Vitamin K, Arsenic, Boron, Chromium, Copper, Iodine, Iron, Manganese, Molybdenum, Nickel, Silicon, Vanadium, and Zinc (http:/ / books. nap. edu/ openbook. php?record_id=10026& page=258). National Academies Press. pp. 258–259. . [46] " Sources of iodine (http:/ / www. iccidd. org/ pages/ iodine-deficiency/ sources-of-iodine. php)". International Council for the Control of Iodine Deficiency Disorders. . [47] " MedlinePlus Medical Encyclopedia: Iodine in diet (http:/ / www. nlm. nih. gov/ medlineplus/ ency/ article/ 002421. htm)". . [48] Wu T, Liu GJ, Li P, Clar C (2002). "Iodised salt for preventing iodine deficiency disorders". Cochrane Database Syst Rev (3): CD003204. doi: 10.1002/14651858.CD003204 (http:/ / dx. doi. org/ 10. 1002/ 14651858. CD003204). PMID 12137681. [49] Felig, Philip; Frohman, Lawrence A. (2001). " Endemic Goiter (http:/ / books. google. de/ books?id=AZUUGrp6yUgC& pg=RA1-PA351)". Endocrinology & metabolism. McGraw-Hill Professional. ISBN 9780070220010. . [50] " Micronutrients - Iodine, Iron and Vitamin A (http:/ / www. unicef. org/ nutrition/ index_iodine. html)". UNICEF. . [51] " Frequently Asked Questions on Potassium Iodide (http:/ / www. fda. gov/ Drugs/ EmergencyPreparedness/ BioterrorismandDrugPreparedness/ ucm072265. htm)". Food and Drug Administration. . Retrieved 2009-06-06. [52] " Potassium Iodide as a Thyroid Blocking Agent in Radiation Emergencies (http:/ / www. thefederalregister. com/ d. p/ 2001-12-11-01-30492)". Food and Drug Administration. . Retrieved 2009-06-06. [53] [54] [55] [56] [57] [58] [59]
http:/ / lpi. oregonstate. edu/ infocenter/ minerals/ iodine/ http:/ / www. atsdr. cdc. gov/ csem/ iodine/ http:/ / chemicalelements. com/ elements/ i. html http:/ / whqlibdoc. who. int/ publications/ 2004/ 9241592001. pdf http:/ / www. iodinenetwork. net/ http:/ / www. organic-chemistry. org/ chemicals/ oxidations/ iodine. shtm http:/ / www. webelements. com/ webelements/ elements/ text/ I/ index. html
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16
Xenon
1
Xenon iodine ← xenon → caesiumKr ↑ Xe ↓ Rn
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
54Xe Periodic table
Appearance colorless gas General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensityMelting pointBoiling pointTriple pointCritical pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
83
92
103
117
137
165
Atomic properties Oxidation states ElectronegativityIonization energies 2nd: 2046.4 kJ·mol−1 3rd: 3099.4 kJ·mol−1Covalent radiusVan der Waals radius Miscellaneous Crystal structureMagnetic orderingThermal conductivitySpeed of soundCAS registry number Most stable isotopes Main article: Isotopes of xenon iso
N.A.
half-life
124
0.095%
124
125
syn
16.9 h
126
0.089%
126
127
syn
36.345 d
Xe Xe Xe Xe
DM
DE (MeV)
DP
Xe is stable with 70 neutron ε
1.652
125
0.662
127
I
Xe is stable with 72 neutron ε
I
Xenon
2
128
1.91%
128
129
26.4%
129
130
4.07%
130
131
21.2%
131
132
26.9%
132
133
syn
5.247 d
134
10.4%
134
135
syn
9.14 h
136
8.86%
136
Xe Xe Xe Xe Xe Xe Xe Xe Xe
Xe is stable with 74 neutron Xe is stable with 75 neutron Xe is stable with 76 neutron Xe is stable with 77 neutron Xe is stable with 78 neutron β−
0.427
133
1.16
135
Cs
Xe is stable with 80 neutron β−
Cs
Xe is stable with 82 neutron
xenon, Xe, 54 noble gases 18, 5, p131.293(6) g·mol−1 [Kr] 5s2 4d10 5p6 2, 8, 18, 18, 8 (Image) gas (0 °C, 101.325 kPa) 5.894 g/L (101.325 kPa) 161.4 K,−111.7 °C,−169.1 °F (101.325 kPa) 165.03 K,−108.12 °C,−162.62 °F 161.405 K (-112°C), 81.6[1] kPa 289.77 K, 5.841 MPa (101.325 kPa) 2.27 kJ·mol−1 (101.325 kPa) 12.64 kJ·mol−1 (100 kPa, 25 °C) 20.786 J·mol−1·K−10, +1, +2, +4, +6, +8 (rarely more than 0) (weakly acidic oxide) 2.6 (Pauling scale) 1st: 1170.4 kJ·mol−1140±9 pm 216 pm face-centered cubic diamagnetic[2] (300 K) 5.65x10-3 W·m−1·K−1 (liquid) 1090 m/s; (gas) 169 m/s 7440-63-3 Xenon (pronounced /ˈzɛnɒn/[3] or
[4]
English pronunciation: /ˈziːnɒn/
)
is a chemical element represented by the symbol Xe. Its atomic number is 54. A colorless, heavy, odorless noble gas, xenon occurs in the Earth's atmosphere in trace amounts.[5] Although generally unreactive, xenon can undergo a few chemical reactions such as the formation of xenon hexafluoroplatinate, the first noble gas compound to be synthesized.[6] [7] [8] Naturally occurring xenon consists of nine stable isotopes. There are also over 40 unstable isotopes that undergo radioactive decay. The isotope ratios of xenon are an important tool for studying the early history of the Solar System.[9] Xenon-135 is produced as a result of nuclear fission and acts as a neutron absorber in nuclear reactors.[10]
Xenon flash
Xenon is used in flash lamps[11] and arc lamps,[12] and as a general anesthetic.[13] The first excimer laser design used a xenon dimer molecule (Xe2) as its lasing medium,[14] and the earliest laser designs used xenon flash lamps as pumps.[15] Xenon is also being used to search for hypothetical weakly interacting massive particles[16] and as the propellant for ion thrusters in spacecraft.[17]
Xenon
History Xenon was discovered in England by William Ramsay and Morris Travers on July 12, 1898, shortly after their discovery of the elements krypton and neon. They found it in the residue left over from evaporating components of liquid air.[18] [19] Ramsay suggested the name xenon for this gas from the Greek word ξένον [xenon], neuter singular form of ξένος [xenos], meaning 'foreign(er)', 'strange(r)', or 'guest'.[20] [21] In 1902, Ramsay estimated the proportion of xenon in the Earth's atmosphere as one part in 20 million.[22] During the 1930s, engineer Harold Edgerton began exploring strobe light technology for high speed photography. This led him to the invention of the xenon flash lamp, in which light is generated by sending a brief electrical current through a tube filled with xenon gas. In 1934, Edgerton was able to generate flashes as brief as one microsecond with this method.[11] [23] [24] In 1939 Albert R. Behnke Jr. began exploring the causes of "drunkenness" in deep-sea divers. He tested the effects of varying the breathing mixtures on his subjects, and discovered that this caused the divers to perceive a change in depth. From his results, he deduced that xenon gas could serve as an anesthetic. Although Lazharev, in Russia, apparently studied xenon anesthesia in 1941, the first published report confirming xenon anesthesia was in 1946 by J. H. Lawrence, who experimented on mice. Xenon was first used as a surgical anesthetic in 1951 by Stuart C. Cullen, who successfully operated on two patients.[25] In 1960 physicist John H. Reynolds discovered that certain meteorites contained an isotopic anomaly in the form of an overabundance of xenon-129. He inferred that this was a decay product of radioactive iodine-129. This isotope is produced slowly by cosmic ray spallation and nuclear fission, but is produced in quantity only in supernova explosions. As the half-life of 129I is comparatively short on a cosmological time scale, only 16 million years, this demonstrated that only a short time had passed between the supernova and the time the meteorites had solidified and trapped the 129I. These two events (supernova and solidification of gas cloud) were inferred to have happened during the early history of the Solar System, as the 129I isotope was likely generated before the Solar System was formed, but not long before, and seeded the solar gas cloud with isotopes from a second source. This supernova source may also have caused collapse of the solar gas cloud.[26] [27] Xenon and the other noble gases were for a long time considered to be completely chemically inert and not able to form compounds. However, while teaching at the University of British Columbia, Neil Bartlett discovered that the gas platinum hexafluoride (PtF6) was a powerful oxidizing agent that could oxidize oxygen gas (O2) to form dioxygenyl hexafluoroplatinate (O2+[PtF6]−).[28] Since O2 and xenon have almost the same first ionization potential, Bartlett realized that platinum hexafluoride might also be able to oxidize xenon. On March 23, 1962, he mixed the two gases and produced the first known compound of a noble gas, xenon hexafluoroplatinate.[29] [8] Bartlett thought its composition to be Xe+[PtF6]−, although later work has revealed that it was probably a mixture of various xenon-containing salts.[30] [31] [32] Since then, many other xenon compounds have been discovered,[33] along with some compounds of the noble gases argon, krypton, and radon, including argon fluorohydride (HArF),[34] krypton difluoride (KrF2),[35] [36] and radon fluoride.[37] By 1971, more than 80 xenon compounds were known.[38] [39]
3
Xenon
Occurrence and production Xenon is a trace gas in Earth's atmosphere, occurring at 0.087±0.001 parts per million (μL/L), or approximately 1 part per 11.5 million,[40] and is also found in gases emitted from some mineral springs. Some radioactive species of xenon, for example, 133Xe and 135Xe, are produced by neutron irradiation of fissionable material within nuclear reactors.[6] Xenon is obtained commercially as a byproduct of the separation of air into oxygen and nitrogen. After this separation, generally performed by fractional distillation in a double-column plant, the liquid oxygen produced will contain small quantities of krypton and xenon. By additional fractional distillation steps, the liquid oxygen may be enriched to contain 0.1–0.2% of a krypton/xenon mixture, which is extracted either via adsorption onto silica gel or by distillation. Finally, the krypton/xenon mixture may be separated into krypton and xenon via distillation.[41] [42] Extraction of a liter of xenon from the atmosphere requires 220 watt-hours of energy.[43] Worldwide production of xenon in 1998 was estimated at 5,000–7,000 m3.[44] Due to its low abundance, xenon is much more expensive than the lighter noble gases—approximate prices for the purchase of small quantities in Europe in 1999 were 10 €/L for xenon, 1 €/L for krypton, and 0.20 €/L for neon.[44] Xenon is relatively rare in the Sun's atmosphere, on Earth, and in asteroids and comets. The atmosphere of Mars shows a xenon abundance similar to that of Earth: 0.08 parts per million,[45] however Mars shows a higher proportion of 129Xe than the Earth or the Sun. As this isotope is generated by radioactive decay, the result may indicate that Mars lost most of its primordial atmosphere, possibly within the first 100 million years after the planet was formed.[46] [47] By contrast, the planet Jupiter has an unusually high abundance of xenon in its atmosphere; about 2.6 times as much as the Sun.[48] This high abundance remains unexplained and may have been caused by an early and rapid buildup of planetesimals—small, subplanetary bodies—before the presolar disk began to heat up.[49] (Otherwise, xenon would not have been trapped in the planetesimal ices.) Within the Solar System, the nucleon fraction for all isotopes of xenon is 1.56 × 10−8, or one part in 64 million of the total mass.[50] The problem of the low terrestrial xenon may potentially be explained by covalent bonding of xenon to oxygen within quartz, hence reducing the outgassing of xenon into the atmosphere.[51] Unlike the lower mass noble gases, the normal stellar nucleosynthesis process inside a star does not form xenon. Elements more massive than iron-56 have a net energy cost to produce through fusion, so there is no energy gain for a star to create xenon.[52] Instead, many isotopes of xenon are formed during supernova explosions.[53]
Characteristics An atom of xenon is defined as having a nucleus with 54 protons. At standard temperature and pressure, pure xenon gas has a density of 5.761 kg/m3, about 4.5 times the surface density of the Earth's atmosphere, 1.217 kg/m3.[54] As a liquid, xenon has a density of up to 3.100 g/mL, with the density maximum occurring at the triple point.[55] Under the same conditions, the density of solid xenon, 3.640 g/cm3, is larger than the average density of granite, 2.75 g/cm3.[55] Using gigapascals of pressure, xenon has been forced into a metallic phase.[56] Solid xenon changes from face-centered cubic (FCC) to hexagonal close packed (HCP) crystal phase under pressure and begins to turn metallic at about 140 GPa, with no
4
Xenon
5
noticeable volume change in the HCP phase. It is completely metallic at 155 GPa. When metalized, xenon looks sky blue because an indirect bandgap transition (in the electronic structure) allows it to absorb red light as had been previously calculated.[57] [58] Metallic xenon is also transparent to visible light. The possibility that metallic xenon could react with iron under pressure deep in the earth was considered a possible reason for xenon's rarity in the atmosphere. However, experimental data and calculations have disproved this possibility. Xenon is a member of the zero-valence elements that are called noble or inert gases. It is inert to most common chemical reactions (such as combustion, for example) because the outer valence shell contains eight electrons. This produces a stable, minimum energy configuration in which the outer electrons are tightly bound.[59] However, xenon can be oxidized by powerful oxidizing agents, and many xenon compounds have been synthesized. Xenon in shaped Geissler tubes
In a gas-filled tube, xenon emits a blue or lavenderish glow when the gas is excited by electrical discharge. Xenon emits a band of emission lines that span the visual spectrum,[60] but the most intense lines occur in the region of blue light, which produces the coloration.[61]
Isotopes Naturally occurring xenon is made of nine stable isotopes, the most of any element with the exception of tin, which has ten. Xenon and tin are the only elements to have more than seven stable isotopes.[62] The isotopes 124Xe, 134Xe and 136Xe are predicted to undergo double beta decay, but this has never been observed so they are considered to be stable.[63] [64] Besides these stable forms, there are over 40 unstable isotopes that have been studied. 129 Xe is produced by beta decay of 129I, which has a half-life of 16 million years, while 131m Xe, 133Xe, 133mXe, and 135Xe are some of the fission products of both 235U and 239 Pu,[65] and therefore used as indicators of nuclear explosions. The various isotopes of xenon are produced from supernova explosions,[53] red giant stars that have exhausted the hydrogen at their cores and entered the asymptotic giant branch, classical novae explosions[66] and the radioactive decay of elements such as iodine, uranium and plutonium.[65] The artificial isotope 135Xe is of considerable significance in the operation of nuclear fission reactors. 135Xe has a huge cross section for thermal neutrons, 2.6×106 barns,[10] so it acts as a neutron absorber or "poison" that can slow or stop the chain reaction after a period of operation. This was discovered in the earliest nuclear reactors built by the American Manhattan Project for plutonium production. Fortunately the designers had made provisions in the design to increase the reactor's reactivity (the number of neutrons per fission that go on to fission other atoms of nuclear fuel).[67] 135Xe reactor poisoning played a major role in the Chernobyl disaster.[68] Under adverse conditions, relatively high concentrations of radioactive xenon isotopes may be found emanating from nuclear reactors due to the release of fission products from cracked fuel rods,[69] or fissioning of uranium in cooling water.[70]
Xenon
6
Because xenon is a tracer for two parent isotopes, xenon isotope ratios in meteorites are a powerful tool for studying the formation of the solar system. The iodine-xenon method of dating gives the time elapsed between nucleosynthesis and the condensation of a solid object from the solar nebula. Xenon isotopic ratios such as 129Xe/130Xe and 136Xe/130Xe are [9] also a powerful tool for understanding terrestrial differentiation and early outgassing. Excess 129Xe found in carbon dioxide well gases from New Mexico was believed to be from the decay of mantle-derived gases soon after Earth's formation.[65] [71]
Compounds See also: Category:Xenon compounds After the discovery in 1962 by Neil Bartlett that xenon can form chemical compounds, a large number of xenon compounds have been discovered and described. Almost all known xenon compounds contain the electronegative atoms fluorine or oxygen.[72]
Halides Three fluorides are known: XeF , XeF , and XeF . The 2
4
6
fluorides are the starting point for the synthesis of almost all xenon compounds. The difluoride XeF
2
is formed when a mixture of
fluorine and xenon gases are exposed to ultraviolet light.[73] Long-term heating of XeF2 at high temperatures under an NiF2 catalyst yields XeF6.[74] Pyrolysis of XeF6 in the presence of NaF yields high-purity XeF4.[75] The xenon fluorides behave as both fluoride acceptors
Xenon tetrafluoride
and fluoride donors, forming salts that contain such cations as XeF+ and Xe2F+3, and anions such as XeF5−, XeF7−, and XeF82−. The green, paramagnetic Xe+2 is formed by the reduction of XeF2 by xenon gas.[76] XeF2 is also able to form coordination complexes with transition metal ions. Over 30 such complexes have been synthesized and characterized.[74] Whereas the xenon fluorides are well-characterized, the other halides are not known, the only exception being the dichloride, XeCl2. Xenon dichloride is reported to be an endothermic, colorless, crystalline compound that decomposes into the elements at 80°C, formed by the high-frequency irradiation of a mixture of xenon, fluorine, and silicon or carbon tetrachloride.[77] However, doubt has been raised as to whether XeCl2 is a real compound and not merely a van der Waals XeF4 crystals, 1962 molecule consisting of weakly bound Xe atoms and Cl2 molecules.[78] Theoretical calculations indicate that the linear molecule XeCl2 is less stable than the van der Waals complex.[79]
Xenon
7
Oxides and oxohalides Only two oxides of xenon are known: xenon trioxide (XeO3) and xenon tetroxide (XeO4), both of which are dangerously explosive and powerful oxidising agents. Xenon dioxide (XeO2) remains elusive – only the XeOO+ cation has been identified by infrared spectroscopy in solid argon.[80] Xenon does not react with oxygen directly; the trioxide is formed by the hydrolysis of XeF6:[81] XeF6 + 3 H2O → XeO3 + 6 HF XeO3 is weakly acidic, dissolving in alkali to form unstable xenate salts containing the HXeO4− anion. These unstable salts easily disproportionate into xenon gas and perxenate salts, containing the XeO64− anion.[82] Barium perxenate, when treated with concentrated sulfuric acid, yields gaseous xenon tetroxide:[77] Ba2XeO6 + 2 H2SO4 → 2 BaSO4 + 2 H2O + XeO4 To prevent decomposition, the xenon tetroxide thus formed is quickly cooled to form a pale-yellow solid. It explodes above −35.9 °C into xenon and oxygen gas. A number of xenon oxyfluorides are known, including XeOF2, XeOF4, XeO2F2, and XeO3F2. XeOF2 is formed by the reaction of OF2 with xenon gas at low temperatures. It may also be obtained by the partial hydrolysis of XeF4. It disproportionates at −20 °C into XeF2 and XeO2F2.[83] XeOF4 is formed by the partial hydrolysis of XeF6,[84] or the reaction of XeF6 with sodium perxenate, Na4XeO6. The latter reaction also produces a small amount of XeO3F2. XeOF4 reacts with CsF to form the XeOF4− anion.[83]
Other compounds Recently, there has been an interest in xenon compounds where xenon is directly bonded to a less electronegative element than fluorine or oxygen, particularly carbon.[85] Electron-withdrawing groups, such as groups with fluorine substitution, are necessary to stabilize these compounds.[82] Numerous such compounds have been characterized, including:[83] [86] • • • • • •
C6F5–Xe+–N≡C–CH3, where C6F5– is the pentafluorophenyl group. [C6F5]2Xe C6F5–Xe–X, where X is CN, F, or Cl. R–C≡C–Xe+, where R is C2F5– or tert-butyl. C6F5–XeF+2 (C6F5Xe)2Cl+
Other compounds containing xenon bonded to a less electronegative element include F–Xe–N(SO2F)2 and F–Xe–BF2. The latter is synthesized from dioxygenyl tetrafluoroborate, O2BF4, at −100 °C.[83] [87] An unusual ion containing xenon is the tetraxenonogold(II) cation, AuXe2+4, which contains Xe–Au bonds.[88] This ion occurs in the compound AuXe4(Sb2F11)2, and is remarkable in having direct chemical bonds between two notoriously unreactive atoms, xenon and gold, with xenon acting as a transition metal ligand. In 1995, M. Räsänen and co-workers, scientists at the University of Helsinki in Finland, announced the preparation of xenon dihydride (HXeH), and later xenon hydride-hydroxide
Xenon (HXeOH), hydroxenoacetylene (HXeCCH), and other Xe-containing molecules.[89] In 2008, Khriachtchev et al. reported the preparation of HXeOXeH by the photolysis of water within a cryogenic xenon matrix.[90] Deuterated molecules, HXeOD and DXeOH, have also been produced.[91]
Clathrates and excimers In addition to compounds where xenon forms a chemical bond, xenon can form clathrates—substances where xenon atoms are trapped by the crystalline lattice of another compound. An example is xenon hydrate (Xe·5.75 H2O), where xenon atoms occupy vacancies in a lattice of water molecules.[92] This clathrate has a melting point of 24 °C.[93] The deuterated version of this hydrate has also been produced.[94] Such clathrate hydrates can occur naturally under conditions of high pressure, such as in Lake Vostok underneath the Antarctic ice sheet.[95] Clathrate formation can be used to fractionally distill xenon, argon and krypton.[96] Xenon can also form endohedral fullerene compounds, where a xenon atom is trapped inside a fullerene molecule. The xenon atom trapped in the fullerene can be monitored via 129 Xe nuclear magnetic resonance spectroscopy. Using this technique, chemical reactions on the fullerene molecule can be analyzed, due to the sensitivity of the chemical shift of the xenon atom to its environment. However, the xenon atom also has an electronic influence on the reactivity of the fullerene.[97] While xenon atoms are at their ground energy state, they repel each other and will not form a bond. When xenon atoms becomes energized, however, they can form an excimer (excited dimer) until the electrons return to the ground state. This entity is formed because the xenon atom tends to fill its outermost electronic shell, and can briefly do this by adding an electron from a neighboring xenon atom. The typical lifetime of a xenon excimer is 1–5 ns, and the decay releases photons with wavelengths of about 150 and 173 nm.[98] [99] Xenon can also form dimers with other elements, such as the halogens bromine, chlorine and fluorine.[100]
Applications Although xenon is rare and relatively expensive to extract from the Earth's atmosphere, it still has a number of applications.
Illumination and optics Gas-discharge lamps Xenon is used in light-emitting devices called xenon flash lamps, which are used in photographic flashes and stroboscopic lamps;[11] to excite the active medium in lasers which then generate coherent light;[101] and, occasionally, in bactericidal lamps.[102] The first solid-state laser, invented in 1960, was pumped by a xenon flash lamp,[15] and lasers used to power inertial confinement fusion are also pumped by xenon flash lamps.[103]
8
Xenon
Continuous, short-arc, high pressure xenon arc lamps have a color temperature closely approximating noon sunlight and are used in solar simulators. That is, the chromaticity of these lamps closely approximates a heated black body radiator that has a temperature close to that observed from the Sun. After they were first introduced during the 1940s, these lamps began replacing the shorter-lived carbon arc lamps in movie projectors.[12] They are employed in typical 35mm and Xenon short-arc lamp IMAX film projection systems, automotive HID headlights and other specialized uses. These arc lamps are an excellent source of short wavelength ultraviolet radiation and they have intense emissions in the near infrared, which is used in some night vision systems. The individual cells in a plasma display use a mixture of xenon and neon that is converted into a plasma using electrodes. The interaction of this plasma with the electrodes generates ultraviolet photons, which then excite the phosphor coating on the front of the display.[104] [105]
Xenon is used as a "starter gas" in high pressure sodium lamps. It has the lowest thermal conductivity and lowest ionization potential of all the non-radioactive noble gases. As a noble gas, it does not interfere with the chemical reactions occurring in the operating lamp. The low thermal conductivity minimizes thermal losses in the lamp while in the operating state, and the low ionization potential causes the breakdown voltage of the gas to be relatively low in the cold state, which allows the lamp to be more easily started.[106] Lasers In 1962, a group of researchers at Bell Laboratories discovered laser action in xenon,[107] and later found that the laser gain was improved by adding helium to the lasing medium.[108] [109] The first excimer laser used a xenon dimer (Xe2) energized by a beam of electrons to produce stimulated emission at an ultraviolet wavelength of 176 nm.[14] Xenon chloride and xenon fluoride have also been used in excimer (or, more accurately, exciplex) lasers.[110] The xenon chloride excimer laser has been employed, for example, in certain dermatological uses.[111]
Anesthesia Xenon has been used as a general anaesthetic, although it is expensive. Even so, anesthesia machines that can deliver xenon are about to appear on the European market.[112] Two mechanisms for xenon anesthesia have been proposed. The first one involves the inhibition of the calcium ATPase pump—the mechanism cells use to remove calcium (Ca2+)—in the cell membrane of synapses.[113] This results from a conformational change when xenon binds to nonpolar sites inside the protein.[114] The second mechanism focuses on the non-specific interactions between the anesthetic and the lipid membrane.[115] Xenon has a minimum alveolar concentration (MAC) of 71%, making it 50% more potent than N2O as an anesthetic.[13] Thus it can be used in concentrations with oxygen that have a lower risk of hypoxia. Unlike nitrous oxide (N2O), xenon is not a greenhouse gas and so it is also viewed as environmentally friendly. Because of the high cost of xenon, however, economic application will require a closed system so that the gas can be recycled, with the
9
Xenon gas being appropriately filtered for contaminants between uses.[43]
Medical imaging Gamma emission from the radioisotope 133Xe of xenon can be used to image the heart, lungs, and brain, for example, by means of single photon emission computed tomography. 133 Xe has also been used to measure blood flow.[116] [117] [118] Nuclei of two of the stable isotopes of xenon, 129Xe and 131Xe, have non-zero intrinsic angular momenta (nuclear spins). When mixed with alkali vapor and nitrogen and exposed to a laser beam of circularly polarized light that is tuned to an absorption line of the alkali atoms, their nuclear spins can be aligned by a spin exchange process in which the alkali valence electrons are spin-polarized by the light and then transfer their polarization to the xenon nuclei via magnetic hyperfine coupling.[119] Typically, pure rubidium metal, heated above 100 °C, is used to produce the alkali vapor. The resulting spin polarization of xenon nuclei can surpass 50% of its maximum possible value, greatly exceeding the equilibrium value dictated by the Boltzmann distribution (typically 0.001% of the maximum value at room temperature, even in the strongest magnets). Such non-equilibrium alignment of spins is a temporary condition, and is called hyperpolarization. Because a 129Xe nucleus has a spin of 1/2, and therefore a zero electric quadrupole moment, the 129Xe nucleus does not experience any quadrupolar interactions during collisions with other atoms, and thus its hyperpolarization can be maintained for long periods of time even after the laser beam has been turned off and the alkali vapor removed by condensation on a room-temperature surface. The time it takes for a collection of spins to return to their equilibrium (Boltzmann) polarization is called the T1 relaxation time. For 129 Xe it can range from several seconds for xenon atoms dissolved in blood[120] to several hours in the gas phase[121] and several days in deeply frozen solid xenon.[122] In contrast, 131 Xe has a nuclear spin value of 3/2 and a nonzero quadrupole moment, and has T1 relaxation times in the millisecond and second ranges.[123] Hyperpolarization renders 129Xe much more detectable via magnetic resonance imaging and has been used for studies of the lungs and other tissues. It can be used, for example, to trace the flow of gases within the [124] [125] lungs.
Other In nuclear energy applications, xenon is used in bubble chambers,[126] probes, and in other areas where a high molecular weight and inert nature is desirable. A by-product of nuclear weapon testing is the release of radioactive Xe-133 and Xe-135. The detection of these isotopes is used to monitor compliance with test ban treaties,[127] as well as to confirm nuclear test explosions by states such as North Korea.[128]
10
Xenon
Liquid xenon is being used as a medium for detecting hypothetical weakly interacting massive particles, or WIMPs. When a WIMP collides with a xenon nucleus, it should, theoretically, strip an electron and create a primary scintillation. By using xenon, this burst of energy could then be readily distinguished from similar events caused by particles such as cosmic rays.[16] However, the XENON experiment at the Gran Sasso National Laboratory in Italy has thus far failed to find any confirmed WIMPs. Even if no WIMPs are detected, the experiment will serve to constrain the properties of A prototype of a xenon ion engine dark matter and some physics models.[129] The current being tested at NASA's Jet Propulsion detector at this facility has demonstrated sensitivity Laboratory. comparable to that of the best cryogenic detectors, and the sensitivity will be increased by an order of magnitude in 2009.[130] Xenon is the preferred fuel for ion propulsion of spacecraft because of its low ionization potential per atomic weight, and its ability to be stored as a liquid at near room temperature (under high pressure) yet be easily converted back into a gas to fuel the engine. The inert nature of xenon makes it environmentally friendly and less corrosive to an ion engine than other fuels such as mercury or caesium. Xenon was first used for satellite ion engines during the 1970s.[131] It was later employed as a propellant for Europe's SMART-1 spacecraft[17] and for the three ion propulsion engines on NASA's Dawn Spacecraft.[132] Chemically, the perxenate compounds are used as oxidizing agents in analytical chemistry. Xenon difluoride is used as an etchant for silicon, particularly in the production of microelectromechanical systems (MEMS).[133] The anticancer drug 5-fluorouracil can be produced by reacting xenon difluoride with uracil.[134] Xenon is also used in protein crystallography. Applied at pressures from 0.5 to 5 MPa (5 to 50 atm) to a protein crystal, xenon atoms bind in predominantly hydrophobic cavities, often creating a high quality, isomorphous, heavy-atom derivative, which can be used for solving the phase problem.[135] [136]
Precautions Xenon gas can be safely kept in normal sealed glass or metal containers at standard temperature and pressure. However, it readily dissolves in most plastics and rubber, and will gradually escape from a container sealed with such materials.[137] Xenon is non-toxic, although it does dissolve in blood and belongs to a select group of substances that penetrate the blood-brain barrier, causing mild to full surgical anesthesia when inhaled in high concentrations with oxygen (see anesthesia subsection above). Many xenon compounds are explosive and toxic due to their strong oxidative properties.[138] At 169 m/s, the speed of sound in xenon gas is slower than that in air[139] (due to the slower average speed of the heavy xenon atoms compared to nitrogen and oxygen molecules), so xenon lowers the resonant frequencies of the vocal tract when inhaled. This produces a characteristic lowered voice pitch, opposite the high-pitched voice caused by inhalation of helium. Like helium, xenon does not satisfy the body's need for oxygen and is a simple asphyxiant; consequently, many universities no longer allow the voice stunt as a general
11
Xenon
12
chemistry demonstration. As xenon is expensive, the gas sulfur hexafluoride, which is similar to xenon in molecular weight (146 versus 131), is generally used in this stunt, although it too is an asphyxiant.[140] It is possible to safely breathe heavy gases such as xenon or sulfur hexafluoride when they include a 20% mixture of oxygen (although xenon at this concentration would be expected to produce the unconsciousness of general anesthesia). The lungs mix the gases very effectively and rapidly, so that the heavy gases are purged along with the oxygen and do not accumulate at the bottom of the lungs.[141] There is, however, a danger associated with any heavy gas in large quantities: it may sit invisibly in a container, and if a person enters a container filled with an odorless, colorless gas, they may find themselves breathing it unknowingly. Xenon is rarely used in large enough quantities for this to be a concern, though the potential for danger exists any time a tank or container of xenon is kept in an unventilated space.[142]
See also • Penning mixture
External links • • • •
WebElements.com – Xenon [143] USGS Periodic Table - Xenon [144] EnvironmentalChemistry.com - Xenon Xenon as an anesthetic [146]
[145]
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Xenon [103] Skeldon, M.D. (1997). " Thermal distortions in laser-diode- and flash-lamp-pumped Nd:YLF laser rods (http:/ / www. lle. rochester. edu/ pub/ review/ v71/ 6_thermal. pdf)" (PDF). LLE Review 71: 137–144. . Retrieved 2007-02-04. [104] Anonymous. " The plasma behind the plasma TV screen (http:/ / www. plasmatvscience. org/ theinnerworkings. html)". Plasma TV Science. . Retrieved 2007-10-14. [105] Marin, Rick (March 21, 2001). " Plasma TV: That New Object Of Desire (http:/ / www. nytimes. com/ 2001/ 03/ 25/ style/ plasma-tv-that-new-object-of-desire. html?sec=& spon=)". The New York Times. . Retrieved 2009-04-03. [106] Waymouth, John (1971). Electric Discharge Lamps. Cambridge, MA: The M.I.T. Press. ISBN 0262230488. [107] Patel, C. K. N.; Bennett Jr., W. R.; Faust, W. L.; McFarlane, R. A. (August 1, 1962). " Infrared spectroscopy using stimulated emission techniques (http:/ / prola. aps. org/ abstract/ PRL/ v9/ i3/ p102_1)". Physical Review Letters 9 (3): 102–104. doi: 10.1103/PhysRevLett.9.102 (http:/ / dx. doi. org/ 10. 1103/ PhysRevLett. 9. 102). . [108] Patel, C. K. N.; Faust, W. L.; McFarlane, R. A. (December 1, 1962). "High gain gaseous (Xe-He) optical masers". Applied Physics Letters 1 (4): 84–85. doi: 10.1063/1.1753707 (http:/ / dx. doi. org/ 10. 1063/ 1. 1753707). [109] Bennett, Jr., W. R. (1962). "Gaseous optical masers". Applied Optics Supplement 1: 24–61. [110] " Laser Output (http:/ / www. rstp. uwaterloo. ca/ laser/ documents/ laser_types. html)". University of Waterloo. . Retrieved 2007-10-07. [111] Baltás, E.; Csoma, Z.; Bodai, L.; Ignácz, F.; Dobozy, A.; Kemény, L. (July 2006). "Treatment of atopic dermatitis with the xenon chloride excimer laser". Journal of the European Academy of Dermatology and Venereology 20 (6): 657–660. doi: 10.1111/j.1468-3083.2006.01495.x (http:/ / dx. doi. org/ 10. 1111/ j. 1468-3083. 2006. 01495. x). [112] Tonner, P. H. (2006). "Xenon: one small step for anaesthesia…? (editorial review)". Current Opinion in Anaesthesiology 19 (4): 382–384. doi: 10.1097/01.aco.0000236136.85356.13 (http:/ / dx. doi. org/ 10. 1097/ 01. aco. 0000236136. 85356. 13). [113] Franks first=John J.; Horn, Jean-Louis; Janicki, Piotr K.; Singh, Gurkeerat (1995). "Halothane, Isoflurane, Xenon, and Nitrous Oxide Inhibit Calcium ATPase Pump Activity in Rat Brain Synaptic Plasma Membranes.". Anesthesiology 82 (1): 108–117. doi: 10.1097/00000542-199501000-00015 (http:/ / dx. doi. org/ 10. 1097/ 00000542-199501000-00015). [114] Lopez, Maria M.; Kosk-Kosicka, Danuta (1995). "How do volatile anesthetics inhibit Ca2+-ATPases?". Journal of Biological Chemistry 270 (47): 28239–28245. doi: 10.1074/jbc.270.47.28239 (http:/ / dx. doi. org/ 10. 1074/ jbc. 270. 47. 28239). PMID 7499320. [115] Heimburg, T.; Jackson A. D. (2007). "The thermodynamics of general anesthesia". Biophysical Journal 92 (9): 3159–65. doi: 10.1529/biophysj.106.099754 (http:/ / dx. doi. org/ 10. 1529/ biophysj. 106. 099754). PMID 17293400. [116] Van Der Wall, Ernst (1992). What's New in Cardiac Imaging?: SPECT, PET, and MRI. Springer. ISBN 0792316150. [117] Frank, John (1999). " Introduction to imaging: The chest (http:/ / student. bmj. com/ issues/ 04/ 01/ education/ 8. php)". Student BMJ 12: 1–44. . Retrieved 2008-06-04. [118] Chandak, Puneet K. (July 20, 1995). " Brain SPECT: Xenon-133 (http:/ / brighamrad. harvard. edu/ education/ online/ BrainSPECT/ Theory/ Xenon133. html)". Brigham RAD. . Retrieved 2008-06-04. [119] Otten, Ernst W. (2004). "Take a breath of polarized noble gas". Europhysics News 35 (1): 16. doi: 10.1051/epn:2004109 (http:/ / dx. doi. org/ 10. 1051/ epn:2004109). [120] Wolber, J.; Cherubini, A.; Leach, M. O.; Bifone, A. (2000). "On the oxygenation-dependent 129Xe T1 in blood". NMR in Biomedicine 13 (4): 234-237. doi: 10.1002/1099-1492(200006)13:4%3C234::AID-NBM632%3E3.0.CO;2-K (http:/ / dx. doi. org/ 10. 1002/ 1099-1492(200006)13:4%3C234::AID-NBM632%3E3. 0. CO;2-K). [121] Chann, B.; Nelson, I. A.; Anderson, L. W.; Driehuys, B.; Walker, T. G. (2002). "129Xe-Xe molecular spin relaxation". Physical Review Letters 88 (11): 113–201. doi: 10.1103/PhysRevLett.88.113201 (http:/ / dx. doi. org/ 10. 1103/ PhysRevLett. 88. 113201). [122] von Schulthess, Gustav Konrad; Smith, Hans-Jørgen; Pettersson, Holger; Allison, David John (1998). The Encyclopaedia of Medical Imaging. Taylor & Francis. pp. 194. ISBN 1901865134. [123] Warren, W. W.; Norberg, R. E. (1966). "Nuclear Quadrupole Relaxation and Chemical Shift of Xe131 in Liquid and Solid Xenon". Physical Review 148 (1): 402–412. doi: 10.1103/PhysRev.148.402 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 148. 402). [124] Albert, M. S.; Balamore, D. (1998). " Development of hyperpolarized noble gas MRI (http:/ / www. ncbi. nlm. nih. gov/ sites/ entrez?cmd=Retrieve& db=PubMed& list_uids=11543065& dopt=AbstractPlus)". Nuclear Instruments and Methods in Physics Research A 402: 441–453. doi: 10.1016/S0168-9002(97)00888-7 (http:/ / dx. doi. org/ 10. 1016/ S0168-9002(97)00888-7). . Retrieved 2007-10-01.
17
Xenon
18
[125] Irion, Robert (March 23, 1999). " Head Full of Xenon? (http:/ / sciencenow. sciencemag. org/ cgi/ content/ full/ 1999/ 323/ 3)". Science News. . Retrieved 2007-10-08. [126] Galison, Peter Louis (1997). Image and Logic: A Material Culture of Microphysics. University of Chicago Press. pp. 339. ISBN 0226279170. [127] Fontaine, J.-P.; Pointurier, F.; Blanchard, X.; Taffary, T. (2004). "Atmospheric xenon radioactive isotope monitoring". Journal of Environmental Radioactivity 72 (1–2;): 129–135. doi: 10.1016/S0265-931X(03)00194-2 (http:/ / dx. doi. org/ 10. 1016/ S0265-931X(03)00194-2). [128] Garwin, Richard L.; von Hippel Frank N. (November 2006). " A Technical Analysis: Deconstructing North Korea's October 9 Nuclear Test (http:/ / www. armscontrol. org/ act/ 2006_11/ tech)". Arms Control Today (Arms Control Association) 38 (9). . Retrieved 2009-03-26. [129] Schumann, Marc (October 10, 2007). " XENON announced new best limits on Dark Matter (http:/ / xenon. physics. rice. edu/ )". Rice University. . Retrieved 2007-10-08. [130] Boyd, Jade (August 23, 2007). " Rice physicists go deep for 'dark matter' (http:/ / www. media. rice. edu/ media/ NewsBot. asp?MODE=VIEW& ID=9902& SnID=1256234278)". Hubble News Desk. . Retrieved 2007-10-08. [131] Zona, Kathleen (March 17, 2006). " Innovative Engines: Glenn Ion Propulsion Research Tames the Challenges of 21st century Space Travel (http:/ / www. nasa. gov/ centers/ glenn/ about/ fs08grc. html)". NASA. . Retrieved 2007-10-04. [132] " Dawn Launch: Mission to Vesta and Ceres (http:/ / www. jpl. nasa. gov/ news/ press_kits/ dawn-launch. pdf)" (PDF). NASA. . Retrieved 2007-10-01. [133] Brazzle, J. D.; Dokmeci, M. R.; Mastrangelo, C. H. (July 28-August 1, 1975). "Modeling and Characterization of Sacrificial Polysilicon Etching Using Vapor-Phase Xenon Difluoride". Proceedings 17th IEEE International Conference on Micro Electro Mechanical Systems (MEMS). Maastricht, Netherlands: IEEE. pp. 737-740. ISBN 9780780382657. [134] Staff (2007). " Powerful tool (http:/ / acswebcontent. acs. org/ landmarks/ bartlett/ tool. html)". American Chemical Society. . Retrieved 2007-10-10. [135] Staff (December 21, 2004). " Protein Crystallography: Xenon and Krypton Derivatives for Phasing (http:/ / www. srs. ac. uk/ px/ facilities/ xenon_notes_1. html)". PX. . Retrieved 2007-10-01. [136] Drenth, Jan; Mesters, Jeroen (2007). "The Solution of the Phase Problem by the Isomorphous Replacement Method". Principles of Protein X-Ray Crystallography (3rd ed.). New York: Springer. pp. 123–171. doi: 10.1007/0-387-33746-6_7 (http:/ / dx. doi. org/ 10. 1007/ 0-387-33746-6_7). ISBN 978-0-387-33334-2. [137] LeBlanc, Adrian D.; Johnson, Philip C. (1971). "The handling of xenon-133 in clinical studies". Physics in Medicine and Biology 16 (1): 105-109. doi: 10.1088/0031-9155/16/1/310 (http:/ / dx. doi. org/ 10. 1088/ 0031-9155/ 16/ 1/ 310). [138] Finkel, A. J.; Katz, J. J.; Miller, C. E. (April 1, 1968). " Metabolic and toxicological effects of water-soluble xenon compounds are studied (http:/ / ntrs. nasa. gov/ search. jsp?R=306918& id=2& qs=No=40& Ne=26& N=297+ 140& Ns=PublicationYear%7C0)". NASA. . Retrieved 2007-10-04. [139] 169.44 m/s in xenon (at 0°C and 107 KPa), compared to 344 m/s in air. See: Vacek, V.; Hallewell, G.; Lindsay, S. (2001). "Velocity of sound measurements in gaseous per-fluorocarbons and their mixtures". Fluid Phase Equilibria 185: 305–314. doi: 10.1016/S0378-3812(01)00479-4 (http:/ / dx. doi. org/ 10. 1016/ S0378-3812(01)00479-4). [140] Spangler, Steve (2007). " Anti-Helium - Sulfur Hexafluoride (http:/ / www. stevespanglerscience. com/ experiment/ from-donald-duck-to-barry-white-how-gases-change-your-voice)". Steve Spangler Science. . Retrieved 2007-10-04. [141] Yamaguchi, K.; Soejima, K.; Koda, E.; Sugiyama, N (2001). "Inhaling Gas With Different CT Densities Allows Detection of Abnormalities in the Lung Periphery of Patients With Smoking-Induced COPD". Chest Journal 51: 1907–1916. doi: 10.1378/chest.120.6.1907 (http:/ / dx. doi. org/ 10. 1378/ chest. 120. 6. 1907). PMID 11742921. [142] Staff (August 1, 2007). " Cryogenic and Oxygen Deficiency Hazard Safety (http:/ / www-group. slac. stanford. edu/ esh/ hazardous_substances/ cryogenic/ p_hazards. htm)". Stanford Linear Accelerator Center. . Retrieved 2007-10-10. [143] [144] [145] [146]
http:/ / www. webelements. com/ webelements/ elements/ text/ Xe/ index. html http:/ / wwwrcamnl. wr. usgs. gov/ isoig/ period/ xe_iig. html http:/ / environmentalchemistry. com/ yogi/ periodic/ Xe. html http:/ / www. anaesthetist. com/ anaes/ drugs/ xenon. htm
Article Sources and Contributors
Article Sources and Contributors Xenon Source: http://en.wikipedia.org/w/index.php?oldid=307984214 Contributors: -Jade healer-, ABF, AJim, AThing, Aadgray, Aarchiba, Acegikmo1, Adam11010, Adashiel, Agesworth, Ahoerstemeier, Ajd, Akamad, Alansohn, Aliwalla, Allstarecho, Alphachimp, Amatulic, Amicon, Anclation, Andre Engels, AndreasJS, Andrei.hagiescu, Andrew Kanode, Animum, Antandrus, Anthony Appleyard, Anthony Duff, Arakunem, Archimerged, Arteitle, Arx Fortis, AuburnPilot, Azn king28, Babooche, Baccyak4H, Bact, Badocter, Barnaclese, BassT0theFace, Beetstra, Benbest, Benjah-bmm27, Bergsten, BerserkerBen, Big Brother 1984, BigMitch741, Bkell, BlueEarth, Bobo192, Bobthebuilder32, Bongwarrior, BorgQueen, Borislav.dopudja, Bradkittenbrink, Bratsche, Brichcja, Bryan Derksen, Bucketsofg, CYD, Cacahueten, Can't sleep, clown will eat me, Carnildo, Cazort, Cdpl, Chem-awb, Chethanyk, Chiloa, Chininazu12, Chrislk02, Christian List, ClanCC, Coemgenus, Columbo2, Condem, Conversion script, Cool3, Corpx, Corvus cornix, Css, CzarB, DMacks, Dalziel 86, Danh, DanielCD, DanielDeGraaf, Dar-Ape, Darrien, David Latapie, Dawn Bard, Deglr6328, Denton22, DerHexer, Dgrant, Dicklyon, Dkelley123, Donarreiskoffer, Dougofborg, Doulos Christos, Dreish, Duccio, Eclecticology, Edgar181, Edibobb, EelkeSpaak, Egil, El C, Emperorbma, Eng02019, Epbr123, Eregli bob, Erich gasboy, Erik Zachte, Eutactic, Evand, Felix Wan, Femto, FileMaster, Fish and karate, Flaquito1979, Folypeelarks, Fonzy, Friedfish, Fvasconcellos, Fvw, GTBacchus, Gaius Cornelius, Gakrivas, Galwhaa, Gcsuchemistry, Gene Nygaard, Giftlite, Gogo Dodo, Graham87, Greathl, GregorB, Grendelkhan, Grrahnbahr, Gruber76, Hannibal, Harrymph, Hvn0413, Iago4096, Icairns, Ideyal, Iidiggsz, Iliev, Itub, Ivansanchez, Ixfd64, J.delanoy, JDPhD, JNW, Jaan513, Jaknouse, James086, Jaraalbe, Jason Leach, Jeremiah Mountain, Jj137, Joanjoc, John, John254, JohnCD, Jonathunder, Jorge Stolfi, Jose77, KJBracey, Karn, Kay Dekker, Keilana, Kelovy, Klausok, Kostisl, Kowey, Kozuch, Ksanyi, Kuru, Kurykh, Kwamikagami, Kwsn, Leranedo, Lexor, LibLord, LightSabre, Lightmouse, Lloydpick, LouScheffer, LuigiManiac, Luna Santin, MPerel, MacGyverMagic, Malcolm Farmer, Marquez, Materialscientist, Mav, Maximaximax, Maxis ftw, Meelar, Megaman en m, Mentifisto, Mernen, Mgimpel, Michael Devore, Mike Rosoft, Mild Bill Hiccup, Minesweeper, MisfitToys, Mjkper, Mohsen4465, Moshe Constantine Hassan Al-Silverburg, Mraandyy, MyLegAlta, Nappyrootslistener, NeilN, Nergaal, Neucom, Neurolysis, Nihiltres, Noah Salzman, Notchcode, Olaymaster888, Onevalefan, Opabinia regalis, Ossipewsk, Oxymoron83, PAR, Paperlan, Parhamr, Persian Poet Gal, Peter M Dodge, PeterJeremy, Petri Krohn, Pharaoh of the Wizards, Phidauex, Philip Trueman, Pierre.laloë, Pjacobi, Poolkris, Possum, Pstudier, Qrc, Qxz, RJHall, RMFan1, RP88, RTC, Randalllin, Randomblue, Raphael s, Remember, Rettetast, RexNL, Rich Farmbrough, Rjwilmsi, Robert Merkel, Roberta F., Robophilosopher, Romanm, Rootology, RoyBoy, Ruslik0, ST47, Sam Hocevar, Saperaud, Saseigel, Savant13, Sbharris, ScaldingHotSoup, Schneelocke, Scottalter, Sengkang, Sennard, Sfuerst, SimonKSK, SimonP, Sjö, Sl, Smilesfozwood, Soliloquial, Spacepotato, SpikeMolec, Splarka, Squids and Chips, SteinbDJ, Stephenb, Stepheno, Stifynsemons, Stone, SusanLesch, Syrthiss, Tanthalas39, Tarret, Tavilis, Tetracube, TheodoreB, Thingg, Thricecube, Tim Starling, Tim1988, TimVickers, Tiptoety, Tlslights, Tpbradbury, Trieste, Truthanado, Unschool, Until It Sleeps, V1adis1av, V8rik, Vanished user, Velvetron, VigilancePrime, Visor, Vsmith, Vuo, Vuong Ngan Ha, Walkerma, Watch37264, Wavelength, White Cat, Wiki alf, Wikifreak09, WolfmanSF, Xenonice, Yekrats, Yyy, Zoicon5, Zsinj, ^demon, Александър, 592 anonymous edits
Image Sources, Licenses and Contributors file:cubic-face-centered.svg Source: http://en.wikipedia.org/w/index.php?title=File:Cubic-face-centered.svg License: GNU Free Documentation License Contributors: User:Stannered file:Electron shell 054 Xenon.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_054_Xenon.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 Image:Xenon-flash.gif Source: http://en.wikipedia.org/w/index.php?title=File:Xenon-flash.gif License: GNU Free Documentation License Contributors: User:Gmaxwell Image:XeTube.jpg Source: http://en.wikipedia.org/w/index.php?title=File:XeTube.jpg License: unknown Contributors: User:Pslawinski Image:Xenon-tetrafluoride-3D-vdW.png Source: http://en.wikipedia.org/w/index.php?title=File:Xenon-tetrafluoride-3D-vdW.png License: Public Domain Contributors: User:Benjah-bmm27 Image:Xenon tetrafluoride.gif Source: http://en.wikipedia.org/w/index.php?title=File:Xenon_tetrafluoride.gif License: Public Domain Contributors: ABF, Benjah-bmm27, MichaelMaggs, Moink, Paginazero, Shizhao, Walkerma Image:Xenon short arc 1.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Xenon_short_arc_1.jpg License: unknown Contributors: User:Atlant Image:Xenon ion engine prototype.png Source: http://en.wikipedia.org/w/index.php?title=File:Xenon_ion_engine_prototype.png License: Public Domain Contributors: NASA
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19
Caesium
1
Caesium 55
xenon ← caesium → barium
Rb ↑
Cs ↓
Fr Periodic Table - Extended Periodic Table
General Name, Symbol, Number
caesium, Cs, 55
Element category
alkali metals
Group, Period, Block
1, 6, s
Appearance
silvery gold
Standard atomic weight
132.9054519(2) g·mol
Electron configuration
[Xe] 6s
−1
1
Electrons per shell
2, 8, 18, 18, 8, 1 Physical properties
Phase
solid
Density (near r.t.)
1.93 g·cm−3
Liquid density at m.p.
1.843 g·cm−3
Melting point
301.59 K (28.44 °C, 83.19 °F)
Boiling point
944 K (671 °C, 1240 °F)
Critical point
1938 K, 9.4 MPa
Heat of fusion
2.09 kJ·mol−1
Heat of vaporization
63.9 kJ·mol−1
Specific heat capacity
(25 °C) 32.210 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
418
469
534
623
750
940
Caesium
2 Atomic properties
Crystal structure
body centered cubic
Oxidation states
1 (strongly basic oxide)
Electronegativity
0.79 (Pauling scale)
Ionization energies
1st: 375.7 kJ/mol 2nd: 2234.3 kJ/mol 3rd: 3400 kJ/mol
Atomic radius
265 pm
Covalent radius
244±11 pm Miscellaneous [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(20 °C) 205 n Ω·m
Thermal conductivity
(300 K) 35.9 W·m
Thermal expansion
(25 °C) 97 µm·m
Young's modulus
1.7 GPa
Bulk modulus
1.6 GPa
−1
−1
Mohs hardness
−1
·K
·K
−1
0.2
Brinell hardness
0.14 MPa
CAS registry number
7440-46-2 Most-stable isotopes
Main article: Isotopes of caesium iso 133
Cs
134
Cs
NA 100% syn
half-life
DM
DE (MeV)
DP
133
Cs is stable with 78 neutron
65.159 Ms (2.0648 y)
ε
1.229
134
β−
2.059
134
Xe Ba
135
trace
73 Ts (2.3 My)
β−
0.269
135
137
syn
948.9 Ms (30.07 y)
β−
1.176
137
Cs Cs
Ba Ba
References
Caesium or cesium (pronounced /ˈsiːziəm/) is the chemical element with the symbol Cs and atomic number 55. It is a soft, silvery-gold alkali metal with a melting point of 28 °C (83 °F), which makes it one of only five metals that are liquid at or near room temperature.[2] Caesium is most notably used in atomic clocks. Caesium is the international spelling standardized by the IUPAC, but in the United States it is more commonly spelled as cesium.[3]
Caesium
3
Characteristics The emission spectrum of caesium has two bright lines in the blue area of the spectrum along with several other lines in the red, yellow, and green areas. This metal is silvery gold in color and is both soft and ductile. Caesium has the lowest ionization potential of the chemical elements. Caesium is the least abundant of the five non-radioactive alkali metals. (Francium is the least common alkali metal but it has no stable isotopes.[4] ). Caesium, gallium, francium, rubidium, and mercury are the only pure metals liquid at or near room temperature. (Some sodium-potassium alloys are also liquid at room temperature.) Caesium reacts explosively in cold water and also reacts with ice at temperatures above −116 °C (−177 °F, 157 K). High purity caesium metal.
Caesium hydroxide (CsOH) is a very strong base and will rapidly etch the surface of glass. CsOH is often stated to be the "strongest base", but in fact many compounds such as n-butyllithium and sodium amide are stronger but are not classic hydroxide bases and are destroyed by water.
Applications Probably the most widespread use of caesium today is in caesium formate-based drilling fluids for the oil industry. The high density of the caesium formate brine (up to 2.3 sg), coupled with the relatively benign nature of natural caesium (which has minimal radioactivity because it is almost entirely composed of a stable istotope), reduces the requirement for toxic high-density suspended solids in the drilling fluid, which is a significant technological, engineering and environmental advantage.[5] [6] Caesium is also used in atomic clocks, which are accurate to seconds over many thousands of years. Since 1967, the International System of Measurements has based its unit of time, the second, on the properties of caesium. The International System of Units (SI) defines the second as 9,192,631,770 cycles of the radiation, which corresponds to the transition between two hyperfine energy levels of the ground state of the 133Cs atom. •
134
Cs has been used in hydrology as a measure of caesium output by the nuclear power industry. This isotope is used because, while it is less prevalent than either 133Cs or 137 Cs, 134Cs can be produced solely by nuclear reactions. 135Cs has also been used in this function. • Like other elements of group 1, caesium has a great affinity for oxygen and is used as a "getter" in vacuum tubes. • This metal is also used in photoelectric cells due to its ready emission of electrons. • Caesium was used as a propellant in early ion engines. It used a method of ionization to strip the outer electron from the propellant by simple contact with tungsten. Caesium use as a propellant was discontinued when Hughes Research Laboratory conducted a study finding xenon gas as a suitable replacement. • Caesium is used as a catalyst in the hydrogenation of certain organic compounds.
Caesium • Radioactive isotopes of caesium are used in the medical field to treat certain types of cancer. • Caesium fluoride is widely used in organic chemistry as a base and as a source of anhydrous fluoride ion. • Caesium vapor is used in many common magnetometers. • Because of their high density, caesium chloride solutions are commonly used in molecular biology for density gradient ultracentrifugation, primarily for the isolation of viral particles, subcellular organelles and fractions, and nucleic acids from biological samples. • Caesium nitrate is used as an oxidizer and pyrotechnic colorant to burn silicon in infrared flares[7] such as the LUU-19 flare,[8] because it emits much of its light in the near infrared spectrum. • Caesium-137 is an extremely common radioisotope used as a gamma-emitter in industrial applications such as: • moisture/density gauges • leveling gauges • thickness gauges • well logging devices which are used to measure the electron density, which is analogous to the bulk density, of the rock formations. • Caesium is also used as an internal standard in spectrophotometry. • Caesium has been used to reduce the radar signature of exhaust plumes in military aircraft.
History Caesium (Latin caesius meaning "blueish grey")[9] [10] was spectroscopically discovered by Robert Bunsen and Gustav Kirchhoff in 1860 in mineral water from Dürkheim, Germany. The residues of 44,000 liters of mineral water yielded several grams of caesium salt for further analysis. Its identification was based upon the bright blue lines in its spectrum and it was the first element discovered by spectrum analysis.[11] The first caesium metal was produced in 1882 by electrolysis of caesium chloride by Carl Setterberg. Setterberg received his PhD from Kekule and Bunsen for this work. Historically, the most important use for caesium has been in research and development, primarily in chemical and electrical applications.
4
Caesium
5
Occurrence An alkali metal, caesium occurs in lepidolite, pollucite (hydrated silicate of aluminium and caesium) and within other sources. One of the world's most significant and rich sources of this metal is at Bernic Lake in Manitoba. The deposits there are estimated to contain 300,000 metric tons of pollucite ore at an average of composition of 20% caesium by weight. It can be isolated by electrolysis of fused caesium cyanide and in a number of other ways. Pollucite, a caesium mineral. Exceptionally pure and gas-free caesium can be made by the thermal decomposition of caesium azide. The primary compounds of caesium are caesium chloride and its nitrate. The price of caesium metal in 1997 was about US$30 per gram, but its compounds are much cheaper.
Isotopes Caesium has at least 39 known isotopes, which is more than any other element except francium. The atomic masses of these isotopes range from 112 to 151. Even though this element has a large number of isotopes, it has only one naturally occurring stable isotope, 133 Cs. Most of the other isotopes have half-lives from a few days to fractions of a second. The radiogenic isotope 137Cs has been used in hydrologic studies, analogous to the use of 3 H. 137Cs is produced from the detonation of nuclear weapons and is produced in nuclear power plants, and was released to the atmosphere most notably from the 1986 Chernobyl meltdown. This isotope (137Cs) is one of the numerous products of fission, directly issued from the fission of uranium. Beginning
in
1945
with
the
commencement of nuclear weapons testing, 137Cs was released into the atmosphere where it is not absorbed readily into solution and is returned to the surface of the earth as a component of radioactive fallout. Once 137Cs enters the ground water, it is deposited on soil surfaces and removed from the landscape primarily by particle transport. As a result, the input function of these isotopes High purity caesium-133 (preserved under argon) cannot be estimated as a function of time. Caesium-137 has a half-life of 30.17 years. It decomposes to barium-137m (a short-lived product of decay) then to a form of nonradioactive barium.
Caesium
Precautions All alkali metals are highly reactive. Caesium, being one of the heavier alkali metals, is also one of the most reactive and is highly explosive when it comes in contact with water. The hydrogen gas produced by the reaction is heated by the thermal energy released at the same time, causing ignition and a violent explosion (the same as all alkali metals) - but caesium is so reactive that this explosive reaction can even be triggered by cold water or ice at temperatures down to -116°C. Caesium is highly pyrophoric and ignites spontaneously in air to form caesium hydroxide and various oxides. Caesium hydroxide is an extremely strong base, and can rapidly corrode glass. Caesium compounds are rarely encountered by most persons. All caesium compounds should be regarded as mildly toxic because of its chemical similarity to potassium. Large amounts cause hyperirritability and spasms, but such amounts would not ordinarily be encountered in natural sources, so Cs is not a major chemical environmental pollutant. Rats fed caesium in place of potassium in their diet die, so this element cannot replace potassium in function in rats. The isotopes 134Cs and 137Cs (present in the biosphere in small amounts as a result of radiation leaks) represent a radioactivity burden which varies depending on location. Radiocaesium does not accumulate in the body as effectively as many other fission products (such as radioiodine and radiostrontium), which are actively accumulated by the body; as with other alkali metals, it washes out of the body relatively quickly in the sweat and urine, unlike strontium which accumulates in the bone.
See also • Caesium-137 • Goiânia accident, a major radioactive contamination incident involving a rod of caesium chloride. • Caesium compounds • Dirty bomb
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] Along with rubidium (39 °C [102 °F]), francium (27 °C [81 °F]), mercury (−39 °C [−38 °F]), and gallium (30 °C [86 °F]). Bromine is also liquid at room temperature (-7.2 °C, 19 °F) but it is not a metal, but a halogen. [3] IUPAC Periodic Table of the Elements (http:/ / old. iupac. org/ reports/ periodic_table/ index. html) [4] Adloff, Jean-Pierre; George B. Kauffman (09/23 2005). " Francium (Atomic Number 87), the Last Discovered Natural Element (http:/ / chemeducator. org/ sbibs/ s0010005/ spapers/ 1050387gk. htm)". The Chemical Educator 10 (5). doi: 10.1333/s00897050956a (http:/ / dx. doi. org/ 10. 1333/ s00897050956a). . Retrieved 2006-05-16. [5] Drilling and Completing Difficult HP/HT Wells With the Aid of Cesium Formate Brines-A Performance Review (http:/ / www. spe. org/ elibinfo/ eLibrary_Papers/ spe/ 2006/ 06DC/ SPE-99068-MS/ SPE-99068-MS. htm) [6] Overview: Cesium Formate Fluids (http:/ / w1. cabot-corp. com/ controller. jsp?N=23+ 4294966885+ 1000& entry=product) [7] United States Patent 6230628: Infrared illumination compositions and articles containing the same (http:/ / www. freepatentsonline. com/ 6230628. html) [8] LUU-19 Flare (http:/ / www. fas. org/ man/ dod-101/ sys/ dumb/ luu19. htm)
6
Caesium [9] Bunsen quotes Aulus Gellius Noctes Atticae II, 26 by Nigidius Figulus: Nostris autem veteribus caesia dicts est quae Graecis, ut Nigidus ait, de colore coeli quasi coelia. [10] Oxford English Dictionary, 2nd Edition [11] G. Kirchhoff, R. Bunsen (1861). "Chemische Analyse durch Spectralbeobachtungen". Annalen der Physik und Chemie 189 (7): 337–381. doi: 10.1002/andp.18611890702 (http:/ / dx. doi. org/ 10. 1002/ andp. 18611890702).
• Los Alamos National Laboratory - Cesium (http:/ / periodic. lanl. gov/ elements/ 55. html) • Daniel A. Steck. " Cesium D Line Data (http:/ / george. ph. utexas. edu/ ~dsteck/ alkalidata/ cesiumnumbers. pdf)". Los Alamos National Laboratory (technical report LA-UR-03-7943). http:/ / george. ph. utexas. edu/ ~dsteck/ alkalidata/ cesiumnumbers. pdf.
External links • WebElements.com – Caesium (http:/ / www. webelements. com/ webelements/ elements/ text/ Cs/ index. html) • Humor site dedicated to caesium (http:/ / www. cs. rochester. edu/ users/ faculty/ nelson/ cesium/ cesium_faq. html) pnb:میزیس
7
Article Sources and Contributors
Article Sources and Contributors Caesium Source: http://en.wikipedia.org/w/index.php?oldid=308638115 Contributors: 203.109.250.xxx, A2Kafir, Afudge, Agateller, Ahoerstemeier, Alan Liefting, Alansohn, Anand Karia, Ancheta Wis, Anclation, Andre Engels, AndreasJS, Andrevan, Antandrus, Anthony Appleyard, Archimerged, Arkuat, Art LaPella, Asmeurer, BGManofID, Baccyak4H, Beetstra, Benbest, BerserkerBen, Billybobjoe, BlueEarth, Blurpeace, Bobo The Ninja, Bobo192, Borislav Dopudja, BrOnXbOmBr21, Bryan Derksen, Bubba hotep, CYD, CanisRufus, CarlOnFire, Carnildo, Catgut, Chris 73, Cleared as filed, Coffee, Colbuckshot, Conversion script, Crazycharles80, Cyrus Andiron, DV8 2XL, Daeroni, Daniel Quinlan, Daniel bg, Darrien, Dave Runger, David Latapie, Dead3y3, Deglr6328, Delldot, DerHexer, Dethme0w, Dman646, Dnn87, Donarreiskoffer, Doodledoo, DrBob, Drini, Dschwen, DuncanHill, ESkog, Eclecticology, Edgar181, Elassint, Elcobbola, Electricitylikesme, Elliott-rhodes, Emperorbma, Epbr123, Ergzay, Eric119, Erik Zachte, Eszett, Everyking, Fanf, Femto, Fifty Percent Normal, Figure, Flewis, Fonzy, Freestyle-69, Gcsuchemistry, Gene Nygaard, Geoffrey, Gouthamyv, Groogle, Hadal, Hak-kâ-ngìn, Hannibal, Hayabusa future, Hello32020, HenryLi, Hgrosser, Hillhead15, Hitssquad, Huey45, Hwvalley, Icairns, Ideyal, Itai, Itub, Ixfd64, JPD, Jaan513, JackLumber, Jake Wartenberg, Jaraalbe, Jdurg, Jh51681, Jimsmith, Jklin, Joanjoc, Joffan, John, Jokem, Jonathunder, Jorge Stolfi, Jose77, KPbIC, Kajasudhakarababu, Karn, Keenan Pepper, Keilana, Keith Edkins, Kilo-Lima, Klosterdev, Knutux, Kotra, Kurykh, Kwamikagami, La goutte de pluie, Lambiam, Lankiveil, Lawrence Cohen, LeaveSleaves, Liastnir, Linas, LorenzoB, Luk, Lysdexia, Mav, Maximaximax, Mcbf111, Mdf, Menuet, Mets501, Michael Slone, Mike Rosoft, Minesweeper, Moeron, Moriori, Mr0t1633, Muke, Myscrnnm, Nergaal, Nihiltres, Nobbie, Nuttycoconut, Octothorn, Oliver Lineham, Pakaran, Parthian Scribe, Pendragonneo, Pgk, Philip Trueman, Physchim62, Piano non troppo, PierreAbbat, Pifactorial, Plexust, Poolkris, Pusher, Pyroseeker, Qwfp, RTC, Remember, RexNL, Richard W.M. Jones, Rifleman 82, Robert Foley, Romanm, Rominandreu, Run!, Rursus, Salsa Shark, Sandahl, Sango123, Saperaud, Sbharris, Schneelocke, Sciurinæ, Scwlong, Sengkang, Severious, Shaddack, Shanedidona, Shantavira, Shniken1, SimonP, Sionus, Sl, Smack, Solipsist, SpeedyGonsales, Spicymanda024, Spykodemon, Squids and Chips, Srtxg, Standard Deviation, Stevertigo, Stifynsemons, Stone, Strait, Stu42, Supaplex, Swedish fusilier, Tagishsimon, TarisWerewolf, Terence, The Anome, The Ronin, The doctor23, TheEgyptian, Thricecube, Titoxd, Tomlillis, Tr-the-maniac, Trevor MacInnis, Trojancowboy, Troy 07, Truthflux, VJDocherty, Varano, Vaughan Pratt, Velvetron, Vsmith, Vuong Ngan Ha, Walkerma, Whicker117, Whosasking, Wimt, Worapon F.S.Boonkerd, Wwoods, Xasodfuih, Yamamoto Ichiro, Yekrats, Yelyos, YukoValis, ZeroP, Орион955, 401 anonymous edits
Image Sources, Licenses and Contributors image:Cs-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Cs-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Paddy, Saperaud Image:Csmetal.jpg.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Csmetal.jpg.jpg License: Creative Commons Attribution-Sharealike 3.0 Contributors: Original uploader was Dnn87 at en.wikipedia Image:Cesium133Cesium.JPG Source: http://en.wikipedia.org/w/index.php?title=File:Cesium133Cesium.JPG License: unknown Contributors: Dnn87 (talk). Original uploader was Dnn87 at en.wikipedia Image:Pollucite(CesiumMineral)USGOV.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Pollucite(CesiumMineral)USGOV.jpg License: Public Domain Contributors: Aushulz, Ra'ike, Saperaud Image:CsCrystals.JPG Source: http://en.wikipedia.org/w/index.php?title=File:CsCrystals.JPG License: unknown Contributors: User:Dnn87
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8
Barium
1
Barium caesium ← barium → lanthanumSr ↑ Ba ↓ Ra
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
56Ba Periodic table
Appearance silvery white
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
911
1038
1185
1388
1686
2170
Barium
2
Atomic properties Oxidation states ElectronegativityIonization energies 2nd: 965.2 kJ·mol−1 3rd: 3600 kJ·mol−1Atomic radiusCovalent radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusMohs hardnessCAS registry number Most stable isotopes Main article: Isotopes of barium iso
N.A.
half-life
130
0.106%
130
132
0.101%
132
133
syn
10.51 y
134
2.417%
134
135
6.592%
135
136
7.854%
136
137
11.23%
137
138
71.7%
138
Ba Ba Ba Ba Ba Ba Ba Ba
DM
DE (MeV)
DP
Ba is stable with 74 neutron Ba is stable with 76 neutron ε
0.517
133
Cs
Ba is stable with 78 neutron Ba is stable with 79 neutron Ba is stable with 80 neutron Ba is stable with 81 neutron Ba is stable with 82 neutron
barium, Ba, 56 alkaline earth metals2, 6, s137.33 g·mol−1 [Xe] 6s2 2, 8, 18, 18, 8, 2 (Image) solid 3.51 g·cm−3 3.338 g·cm−3 1000 K,727 °C,1341 °F 2170 K,1897 °C,3447 °F 7.12 kJ·mol−1 140.3 kJ·mol−1 (25 °C) 28.07 J·mol−1·K−1 2 (strongly basic oxide) 0.89 (Pauling scale) 1st: 502.9 kJ·mol−1222 pm215±11 pm body-centered cubic paramagnetic (20 °C) 332 nΩ·m (300 K) 18.4 W·m−1·K−1 (25 °C) 20.6 µm·m−1·K−1 (20 °C) 1620 m/s 13 GPa 4.9 GPa 9.6 GPa 1.25 7440-39-3 Barium (pronounced /ˈbæriəm/) is a chemical element. It has the symbol Ba, and atomic number 56. Barium is a soft silvery metallic alkaline earth metal. It is never found in nature in its pure form due to its reactivity with air. Its oxide is historically known as baryta but it reacts with water and carbon dioxide and is not found as a mineral. The most common naturally occurring minerals are the very insoluble barium sulfate, BaSO4 (barite), and barium carbonate, BaCO3 (witherite). Benitoite is a rare gem containing barium. Metallic barium has few industrial uses, but has been historically used to scavenge air in vacuum tubes. Barium compounds impart a green color to flames and have been used in fireworks. Barium sulfate is used for its heaviness, insolubility, and X-ray opacity. It is used as an insoluble heavy mud-like paste when drilling oil wells, and in purer form, as an X-ray radiocontrast agent for imaging the human gastrointestinal tract. Soluble barium compounds are poisonous due to release of the soluble barium ion, and have been used as rodenticides. New uses for barium continue to be found: it is an essential ingredient in "high temperature" YBCO superconductors.
Barium
3
Characteristics Physical Barium is a soft and ductile metal. Its simple compounds are notable for their relatively high (for an alkaline earth element) specific gravity. This is true of the most common barium-bearing mineral, its sulfate barite BaSO4, also called 'heavy spar' due to the high density (4.5 g/cm³).
Chemical Barium reacts exothermically with oxygen at room temperature to form barium oxide and peroxide. The reaction is violent if barium is powdered. It also reacts violently with dilute acids, alcohol and water Ba + 2 H2O → Ba(OH)2 + H2 (g) At elevated temperatures, barium combines with chlorine, nitrogen and hydrogen to produce BaCl2, Ba3N2 and BaH2, respectively. Barium reduces oxides, chlorides and sulfides of less reactive metals. For example: Ba + CdO → BaO + Cd Ba + ZnCl2 → BaCl2 + Zn 3 Ba + Al2S3 → 3 BaS + 2 Al When heated with nitrogen and carbon, it forms the cyanide: Ba + N2 + 2 C → Ba(CN)2 Barium combines with several metals, including aluminium, zinc, led and tin, forming intermetallic compounds and alloys.[1]
Isotopes Naturally occurring barium is a mix of seven stable isotopes, the most abundant being 138 Ba (71.7 %). There are twenty-two isotopes known, but most of these are highly radioactive and have half-lives in the several millisecond to several day range. The only notable exceptions are 133Ba which has a half-life of 10.51 years, and 137mBa (2.55 minutes).[2]
History Name barium originates from Greek bary, meaning "heavy". Alchemists in the early Middle Ages knew about some barium minerals. Smooth pebble-like stones of mineral barite found in Bologna, Italy were known as "Bologna stones". After exposed to light they would glow for years that attracted them to witches and alchemists.[3] Carl Scheele identified barite in 1774, but did not isolate barium. Barium was isolated, as ions in solution, in 1808 by Sir Humphry Davy in England. The oxidized barium was at first called barote, by Guyton de Morveau, which was changed by Antoine Lavoisier to baryta, from which "barium" was derived to describe the metal.[3]
Barium
4
Occurrence and production The abundance of barium is 0.0425 % in the Earth's crust and 13 µg/L in sea water. It occurs in the minerals barite (as the sulfate) and witherite (as the carbonate).[1] Large deposits of barite are found in China, Germany, India, Morocco, and in the US.[4]
Barite
Because barium quickly becomes oxidized in air, it is difficult to obtain this metal in its pure form. It is primarily found in and extracted from barite. Because barite is so insoluble, it cannot be used directly for the preparation of other barium compounds. Instead, the ore is heated with carbon to reduce it to barium sulfide:[5] BaSO4 + 2 C → BaS + 2 CO2 The barium sulfide is then hydrolyzed or treated with acids to form other barium compounds, such as the chloride, nitrate, and carbonate. Barium is commercially produced through electrolysis of molten barium chloride (BaCl2):
the
(cathode) Ba2+ + 2 e− → Ba Trend in world production of barite
(anode) 2 Cl− → Cl2 (g) + 2 e−
Barium metal is also obtained by the reduction of barium oxide with finely divided aluminum at temperatures between 1100 and 1200 °C: 4 BaO + 2 Al → BaO·Al2O3 + 3 Ba (g) The barium vapor is cooled by means of a water jacket and condensed into the solid metal. The solid block may be cast into rods or extruded into wires. Being a flammable solid, it is packaged under argon in steel containers or plastic bags. [1]
Applications The most important use of elemental barium is as a scavenger removing last traces of oxygen and other gases in television and other electronic tubes. Besides, an isotope of barium, 133Ba, is routinely used as a standard source in the calibration of gamma-ray detectors in nuclear physics studies.[1] Barium is an important component of YBCO superconductors. An alloy of barium with nickel is used in spark plug wire. Barium oxide is used in a coating for the electrodes of fluorescent lamps, which facilitates the release of electrons.
Amoebiasis as seen in radiograph of barium-filled colon
Barium compounds, and especially barite (BaSO4), are extremely important to the petroleum industry. Besides,
Barium
5
• Barite is used in rubber production and in drilling mud, a weighting agent in drilling new oil wells.[4] • Barium sulfate is used as a radiocontrast agent for X-ray imaging of the digestive system ("barium meals" and "barium enemas").[4] Lithopone, a pigment that contains barium sulfate and zinc sulfide, is a permanent white that has good covering power, and does not darken in when exposed to sulfides.[6] • Barium carbonate is a rat poison and can also be used in making bricks. Unlike the sulfate, the carbonate dissolves in stomach acid, allowing it to be poisonous. Barium carbonate is also used in glassmaking. Being a heavy element, barium increases the refractive index and luster of the glass.[4]
Green barium fireworks
• Barium nitrate and chlorate give green colors in fireworks.[7] • Barium peroxide can be used as a catalyst to start an aluminothermic reaction when welding rail tracks together. It can also be used in green tracer ammunition and as a bleaching agent.[8] • Barium titanate was proposed in 2007 to be used in next generation battery technology for electric cars.[9] • Barium fluoride is used for optics in infrared applications, since it is transparent from about 0.15 to 12 microns.[10] • Barium strontium niobate is used in large scale outdoors Holgraphic displays as a "screen". When projected with blue laser light images are 3-D and very realistic.
Precautions Barium powder is pyrophoric - it can explode in contact with air or oxidizing gases. it is likely to explode when combined with halogenated hydrocarbon solvents. It reacts violently with water. All water or acid soluble barium compounds are extremely poisonous. At low doses, barium acts as a muscle stimulant, while higher doses affect the nervous system, causing cardiac irregularities, tremors, weakness, anxiety, dyspnea and paralysis. This may be due to its ability to block potassium ion channels which are critical to the proper function of the nervous system.[1] Barium sulfate can be taken orally because it is highly insoluble in water, and is eliminated completely from the digestive tract.[1] Unlike other heavy metals, barium does not bioaccumulate.[11] [12] However, inhaled dust containing barium compounds can accumulate in the lungs, causing a benign condition called baritosis.[13] Oxidation occurs very easily and, to remain pure, barium should be kept under a petroleum-based fluid (such as kerosene) or other suitable oxygen-free liquids that exclude air. Barium acetate could lead to death in high doses. Marie Robards poisoned her father with the substance in Texas in 1993. She was tried and convicted in 1996.[14]
Barium
6
External links • WebElements.com – Barium [15] • Elementymology & Elements Multidict
[16]
pnb:میریب
References [1] Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds (http:/ / books. google. com/ books?id=Xqj-TTzkvTEC& pg=PA243). McGraw-Hill. pp. 77–78. ISBN 0070494398. . Retrieved 2009-06-06. [2] David R. Lide, Norman E. Holden (2005). "Section 11, Table of the Isotopes". CRC Handbook of Chemistry and Physics, 85th Edition. Boca Raton, Florida: CRC Press. [3] Robert E. Krebs (2006). The history and use of our earth's chemical elements: a reference guide (http:/ / books. google. com/ books?id=yb9xTj72vNAC& ). Greenwood Publishing Group. p. 80. ISBN 0313334382. . [4] C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304814. [5] " Toxicological Profile for Barium and Barium Compounds. Agency for Toxic Substances and Disease Registry (http:/ / www. atsdr. cdc. gov/ toxprofiles/ tp24. pdf)". CDC. 2007.. . [6] Chris J. Jones, John Thornback (2007). Medicinal applications of coordination chemistry (http:/ / books. google. co. jp/ books?id=uEJHsZWyO-EC& ). Royal Society of Chemistry. p. 102. ISBN 0854045961. . [7] Michael S. Russell, Kurt Svrcula (2008). Chemistry of Fireworks (http:/ / books. google. co. jp/ books?id=yxRyOf8jFeQC& ). Royal Society of Chemistry. p. 110. ISBN 0854041273. . [8] Brent, G. F. (1995). "Surfactant coatings for the stabilization of barium peroxide and lead dioxide in pyrotechnic compositions". Propellants Explosives Pyrotechnics 20: 300. doi: 10.1002/prep.19950200604 (http:/ / dx. doi. org/ 10. 1002/ prep. 19950200604). [9] " Battery Breakthrough? (http:/ / www. technologyreview. com/ Biztech/ 18086/ )". . Retrieved 2009-06-06. [10] " Crystran Ltd. Optical Component Materials (http:/ / www. crystran. co. uk/ products. asp?productid=75)". . Retrieved 2009-06-06. [11] " Toxicity Profiles, Ecological Risk Assessment (http:/ / www. epa. gov/ region5/ superfund/ ecology/ html/ toxprofiles. htm#ba)". . Retrieved 2009-06-06. [12] Moore, J. W. (1991). Inorganic Contaminants of Surface Waters, Research and Monitoring Priorities. New York: Springer-Verlag. [13] Doig AT (February 1976). " Baritosis: a benign pneumoconiosis (http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?tool=pmcentrez& artid=470358)". Thorax 31 (1): 30–9. doi: 10.1136/thx.31.1.30 (http:/ / dx. doi. org/ 10. 1136/ thx. 31. 1. 30). PMID 1257935. [14] " Boyfriend fight preceded Roanoke mom's slaying (http:/ / www. buffalo. edu/ news/ pdf/ October08/ DallanMorningNewsEwingSlaying. pdf)". . Retrieved 2009-06-06. [15] http:/ / www. webelements. com/ webelements/ elements/ text/ Ba/ index. html [16] http:/ / elements. vanderkrogt. net/ elem/ ba. html
Article Sources and Contributors
Article Sources and Contributors Barium Source: http://en.wikipedia.org/w/index.php?oldid=308365013 Contributors: 10014derek, 123awea, 65.68.87.xxx, Abrech, Ahoerstemeier, AlanBarrett, Alansohn, Ale jrb, Andros 1337, Angolob, Animebill, Antandrus, Antonio Lopez, Arcadian, Archimerged, Arkuat, Bayerischermann, Beetstra, Benbest, Blanchardb, BlueEarth, Bobo192, Borislav Dopudja, Brian Huffman, Brockert, Bryan Derksen, C. Ryan Long, CYD, Cacahueten, Caltas, Camw, Carnildo, Chrisd87, Conversion script, D, DanMS, Darrien, David Latapie, DeadEyeArrow, Dlae, Donarreiskoffer, Download, DryaUnda, Dvptl, Dwmyers, Edgar181, Emc2, Emperorbma, Epbr123, Ergzay, Eric119, Erik Zachte, Excirial, Feezo, Femto, FisherQueen, Fivemack, Fonzy, FreplySpang, Gaius Cornelius, Garrettlloydbirch, Gcsuchemistry, Gmaxwell, Greatpatton, Grendelkhan, Hak-kâ-ngìn, Hannibal, Hawkeye2007, Hellbus, Hongooi, HumanFrailty, Iain99, Icairns, Ideyal, ImperatorExercitus, Ixfd64, J0lt C0la, Ja0492, JackofOz, James P Twomey, Jaraalbe, JeLuF, Jeff Wheeler, Jeronimo, Jfbcubed, Joanjoc, John, John Fader, Jose77, Kalamkaar, Katieh5584, Khukri, KnightLago, Ktsquare, Kwamikagami, LMB, Lawrence Cohen, Little Mountain 5, Looxix, LuigiManiac, MPerel, MaNeMeBasat, Marek69, MarkRose, Materialscientist, Mav, Maxamegalon2000, Maximaximax, Mdf, Mentifisto, MightyWarrior, Minesweeper, Mjk2357, Montgomery '39, Mosca, Myasuda, Nagaminerals, NawlinWiki, Neelix, Nergaal, Nihiltres, Numbo3, Pakaran, Peruvianllama, PierreAbbat, Poolkris, Pyrochem, RJFJR, RTC, Rcawsey, Rdsmith4, Rebecca6789, Redux, Remember, Res2216firestar, Rewster, Reza kalani, Rhopkins8, Romanm, Royalguard11, SEJohnston, Sam Hocevar, Saperaud, Sbharris, Sbrockway, Schneelocke, Sengkang, Shafei, Shizane, Sionus, Sl, SnackPack48, SpeedyGonsales, Squids and Chips, Stifynsemons, Stone, Sunborn, Syrthiss, Tagishsimon, Tetracube, The sunder king, Thinboy00, Thricecube, Thryduulf, Tim Starling, Tsogo3, Urhixidur, Useight, VASANTH S.N., Velvetron, Vsmith, Vuo, Walkerma, Warut, Watch37264, Wrenchelle, Yekrats, Yyy, 359 anonymous edits
Image Sources, Licenses and Contributors file:cubic-body-centered.svg Source: http://en.wikipedia.org/w/index.php?title=File:Cubic-body-centered.svg License: GNU Free Documentation License Contributors: User:Stannered file:Electron shell 056 Barium.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_056_Barium.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 file:Barium unter Argon Schutzgas Atmosphäre.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Barium_unter_Argon_Schutzgas_Atmosphäre.jpg License: Public Domain Contributors: Matthias Zepper Image:Barite.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Barite.jpg License: Public Domain Contributors: Andel, Saperaud Image:BariteWorldProductionUSGS.PNG Source: http://en.wikipedia.org/w/index.php?title=File:BariteWorldProductionUSGS.PNG License: Public Domain Contributors: USGS Image:BariumXray.jpg Source: http://en.wikipedia.org/w/index.php?title=File:BariumXray.jpg License: Public Domain Contributors: CDC/ Dr. Mae Melvin; Dr. E. West of Mobile, AL (Public Health Image Library (PHIL)) Image:2006 Fireworks 1.JPG Source: http://en.wikipedia.org/w/index.php?title=File:2006_Fireworks_1.JPG License: unknown Contributors: Joe Anderson
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
7
Lanthanum
1
Lanthanum 57
barium ← lanthanum → cerium
↑
La ↓
Ac Periodic Table - Extended Periodic Table
General Name, Symbol, Number
lanthanum, La, 57
Element category
lanthanides
Group, Period, Block
3, 6, f
Appearance
silvery white
Standard atomic weight
138.90547(7) g·mol
Electron configuration
[Xe] 5d 6s
−1
1
Electrons per shell
2
2, 8, 18, 18, 9, 2 Physical properties
Phase
solid
Density (near r.t.)
6.162 g·cm−3
Liquid density at m.p.
5.94 g·cm−3
Melting point
1193 K (920 °C, 1688 °F)
Boiling point
3737 K (3464 °C, 6267 °F)
Heat of fusion
6.20 kJ·mol−1
Heat of vaporization
402.1 kJ·mol−1
Specific heat capacity
(25 °C) 27.11 J·mol−1·K−1 Vapor pressure (extrapolated)
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
2005
2208
2458
2772
3178
3726
Atomic properties
Lanthanum
2
Crystal structure
hexagonal
Oxidation states
3, 2 (strongly basic oxide)
Electronegativity
1.10 (Pauling scale)
Ionization energies (more)
1st: 538.1 kJ·mol−1 2nd: 1067 kJ·mol−1 3rd: 1850.3 kJ·mol−1
Atomic radius
187 pm
Covalent radius
207±8 pm Miscellaneous [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(r.t.) (α, poly) 615 nΩ·m
Thermal conductivity
(300 K) 13.4 W·m
Thermal expansion
(r.t.) (α, poly) 12.1 µm/(m·K)
Speed of sound (thin rod)
(20 °C) 2475 m/s
Young's modulus
(α form) 36.6 GPa
Shear modulus
(α form) 14.3 GPa
Bulk modulus
(α form) 27.9 GPa
Poisson ratio
(α form) 0.280
Mohs hardness
2.5
Vickers hardness
491 MPa
Brinell hardness
363 MPa
CAS registry number
−1
·K
−1
7439-91-0 Most-stable isotopes
Main article: Isotopes of lanthanum iso 137
La
138
La
139
La
NA syn 0.09%
99.91%
half-life
DM
DE (MeV)
DP
60,000 yrs
ε
0.600
137
1.05×1011yrs
ε
1.737
138
β−
1.044
138
Ba Ba Ce
139
La is stable with 82 neutron References
Lanthanum (pronounced /ˈlænθənəm/) is a chemical element with the symbol La and atomic number 57. Lanthanum is a silvery white metallic element that belongs to group 3 of the periodic table and is a lanthanoid. It is found in some rare-earth minerals, usually in combination with cerium and other rare earth elements. Lanthanum is malleable, ductile, and soft metal, which oxidizes rapidly when exposed to air. It is produced from minerals
Lanthanum monazite and bastnäsite using a complex multistage extraction process. Lanthanum compounds have numerous applications such as catalyst, additives in glass, carbon lighting for studio lighting and projection, ignition element in lighters and torches, electron cathode, scintillator, and others. Lanthanum carbonate was approved as a medication against renal failure.
Properties Physical Lanthanum is soft, malleable, silvery white metal which has hexagonal crystal structure at room temperature. At 310 °C, lanthanum changes to a face-centered cubic structure, and at 865°C into a body-centered cubic structure.[2] Lanthanum easily oxidizes and therefore used as element only for research purpose. For example, single La atoms have been isolated by implanting them into fullerene molecules.[3] If carbon nanotubes are filled with those lanthanum-encapsulated fullerenes and annealed, metallic nanochains of lanthanum are produced inside carbon nanotubes.[4]
Chemical Lanthanum exhibits two oxidation states, +3 and +2, the former being much more stable. For example, LaH3 is more stable than LaH2.[5] Lanthanum burns readily at 150 °C to form lanthanum(III) oxide: 4 La + 3 O2 → 2 La2O3 However, when exposed to moistured air at room temperature, it forms a hydrated oxide with a large volume increase.[5] Lanthanum is quite electropositive and reacts slowly with cold water and quite quickly with hot water to form lanthanum hydroxide: 2 La (s) + 6 H2O (g) → 2 La(OH)3 (aq) + 3 H2 (g) Lanthanum metal reacts with all the halogens. The reaction is vigorous if conducted at above 200 °C: 2 La (s) + 3 F2 (g) → 2 LaF3 (s) 2 La (s) + 3 Cl2 (g) → 2 LaCl3 (s) 2 La (s) + 3 Br2 (g) → 2 LaBr3 (s) 2 La (s) + 3 I2 (g) → 2 LaI3 (s) Lanthanum dissolves readily in dilute sulfuric acid to form solutions containing the La(III) ions, which exist as a [La(OH2)9]3+ complexes:[6] 2 La(s) + 3 H2SO4 (aq) → 2 La3+(aq) + 3 SO2−4 (aq) + 3 H2 (g)
Lanthanum combines with nitrogen, carbon, sulfur, phosphorus, boron, selenium, silicon and arsenic at elevated temperatures, forming binary compounds.[5]
3
Lanthanum
Isotopes Naturally occurring lanthanum is composed of one stable (139La) and one radioactive (138La) isotope, with the stable isotope, 139La, being the most abundant (99.91% natural abundance). 38 radioisotopes have been characterized with the most stable being 138La with a half-life of 1.05×1011 years, and 137La with a half-life of 60,000 years. Most of the remaining radioactive isotopes have half-lives that are less than 24 hours and the majority of these have half lives that are less than 1 minute. This element also has 3 meta states. The isotopes of lanthanum range in atomic weight from 117 u (117La) to 155 u (155La).
History The word lanthanum comes from the Greek λανθανω [lanthanō] = to lie hidden. Lanthanum was discovered in 1839 by Swedish chemist Carl Gustav Mosander, when he partially decomposed a sample of cerium nitrate by heating and treating the resulting salt with dilute nitric acid. From the resulting solution, he isolated a new rare earth he called lantana. Lanthanum was isolated in relatively pure form in 1923.[5] Lanthanum is the most strongly basic of all the trivalent lanthanoids, and this property is what allowed Mosander to isolate and purify the salts of this element. Basicity separation as operated commercially involved the fractional precipitation of the weaker bases (such as didymium) from nitrate solution by the addition of magnesium oxide or dilute ammonia gas. Purified lanthanum remained in solution. (The basicity methods were only suitable for lanthanum purification; didymium could not be efficiently further separated in this manner.) The alternative technique of fractional crystallization was invented by Dmitri Mendeleev, in the form of the double ammonium nitrate tetrahydrate, which he used to separate the less-soluble lanthanum from the more-soluble didymium in the 1870s. This system would be used commercially in lanthanum purification until the development of practical solvent extraction methods that started in the late 1950s. (A detailed process using the double ammonium nitrates to provide 99.99% pure lanthanum, neodymium concentrates and praseodymium concentrates is presented in Callow 1967, at a time when the process was just becoming obsolete.) As operated for lanthanum purification, the double ammonium nitrates were recrystallized from water. When later adapted by Carl Auer von Welsbach for the splitting of didymium, nitric acid was used as solvent to lower the solubility of the system. Lanthanum is relatively easy to purify, since it has only one adjacent lanthanoid, cerium, which itself is very readily removed due to its potential tetravalency. The fractional crystallization purification of lanthanum as the double ammonium nitrate was sufficiently rapid and efficient, that lanthanum purified in this manner was not expensive. The Lindsay Chemical Division of American Potash and Chemical Corporation, for a while the largest producer of rare earths in the world, in a price list dated October 1, 1958 priced 99.9% lanthanum ammonium nitrate (oxide content of 29%) at $3.15 per pound, or $1.93 per pound in 50-pound quantities. The corresponding oxide (slightly purer at 99.99%) was priced at $11.70 or $7.15 per pound for the two quantity ranges. The price for their purest grade of oxide (99.997%) was $21.60 and $13.20, respectively.
4
Lanthanum
Occurrence Although lanthanum belongs to chemical elements group called rare earth metals, it is not rare at all. Lanthanum is available in relatively large quantities (32 ppm in Earth’s crust). "Rare earths" got their name since they were indeed rare as compared to the "common" earths such as lime or magnesia, and historically only a few deposits were known.[5] Monazite (Ce, La, Th, Nd, Y)PO4, and bastnäsite (Ce, La, Y)CO3F, are the principal ores in which lanthanum occurs, in percentages of up to 25 to 38 percent of the Monazite total lanthanoid content. In general, there is more lanthanum in bastnäsite than in monazite. Until 1949, bastnäsite was a rare and obscure mineral, not even remotely contemplated as a potential commercial source for lanthanoids. In that year, the vast deposit at Mountain Pass, California was discovered. This discovery alerted geologists as to the existence of a whole new class of rare earth deposit, the rare-earth bearing carbonatite, other examples of which soon surfaced, particularly in Africa and China.
Production
Lanthanum is most commonly obtained from monazite and bastnäsite. The mineral mixtures are crushed and ground. Monazite, because of its magnetic properties can be separated by repeated electromagnetic separation. After separation, it is treated with hot concentrated sulfuric acid to produce water-soluble sulfates of rare earths. The acidic filtrates are partially neutralized with sodium hydroxide to pH 3-4. Thorium precipitates out of solution as hydroxide and is removed. After that the solution is treated with ammonium oxalate to convert rare earths in to their insoluble oxalates. The oxalates are converted to oxides by annealing. The oxides are dissolved in nitric acid that excludes one of the main components, cerium, whose oxide is insoluble in HNO3. Lanthanum is separated as a double salt with ammonium nitrate by crystallization. This salt is relatively less soluble than other rare earth double salts and therefore stays in the residue.[5] The most efficient separation routine for lanthanum salt from the rare-earth salt solution is however ion exchange. In this process, rare-earth ions are adsorbed onto suitable ion-exchange resin by exchange with hydrogen, ammonium or cupric ions present in the
5
Lanthanum
6
resin. The rare earth ions are then selectively washed out by suitable complexing agent, such as ammonium citrate or nitrilotracetate. Lanthanum can also be separated from solution of rare earth nitrates by liquid-liquid extraction with a suitable organic liquid, such as tributyl phosphalate.[5] Currently, the most widely used extractant for the purification of lanthanum and the other lanthanoids is the 2-ethylhexyl ester of 2-ethylhexylphosphonic acid; this has better handling characteristics than the previously used bis-2-ethylhexyl phosphate. Lanthanum metal is obtained from its oxide by heating it with ammonium chloride or fluoride and hydrofluoric acid at 300-400 °C to produce the chloride or fluoride: La2O3 + 6 NH4Cl → 2 LaCl3 + 6 NH3 + 3 H2O This is followed by reduction with alkali or alkaline earth metals in vacuum or argon atmosphere: LaCl3 + 3 Li → La + 3 LiCl Also pure lanthanum can be produced by electrolysis of molten mixture of anhydrous LaCl3 and NaCl or KCl at elevated temperatures.[5]
Applications The first historical application of lanthanum was in gas lantern mantles. Carl Auer von Welsbach used a mixture of 60% magnesium oxide, 20% lanthanum oxide and 20% yttrium oxide which he called Actinophor, and patented in 1885. The original mantles gave a green-tinted light and were not very successful, and his first company, which established a factory in Atzgersdorf in 1887, failed in 1889.[7] Modern uses of lanthanum include:
A Coleman white gas lantern mantle burning at full brightness.
Lanthanum
7
• Lanthanum oxide and the boride are used in electronic vacuum tubes as hot cathode materials with strong emissivity of electrons. Crystals of LaB6 are used in high brightness, extended life, thermionic electron emission sources for scanning electron microscopes and Hall effect thrusters.[8] • Lanthanum fluoride (LaF3) is an essential component of a heavy fluoride glass named ZBLAN. This glass has superior transmittance in the infrared range and is therefore used for fiber-optical communication systems.[9] • Cerium doped lanthanum bromide and lanthanum chloride are the recent inorganic scintillators which have a combination of high light yield, best energy resolution and fast response. Their high yield converts into superior energy resolution; moreover, the light output is very stable and quite high over a very wide range of temperatures, making it particularly attractive for high temperature applications. These scintillators are already widely used commercially in detectors of neutrons or gamma rays.[10]
LaB6 hot cathode
• Carbon lighting applications, especially by the motion picture industry for studio lighting and projection. These applications consume about 25% of the rare-earth compounds produced.[2] • La O improves the alkali resistance of glass, and is 2
3
Comparison of infrared transmittance of ZBLAN glass and silica
used in making special optical glasses, such as infrared-absorbing glass, as well as camera and telescope lenses, because of the high refractive index and low dispersion of rare-earth glasses.[2] Lanthanum oxide is also used as a grain growth additive during the liquid phase sintering of silicon nitride and zirconium diboride.[11]
• Small amounts of lanthanum added to steel improves its malleability, resistance to impact and ductility. Whereas addition of lanthanum to molybdenum decreases its hardness and sensitivity to temperature variations.[2] • Small amounts of lanthanum are present in many pool products to remove the phosphates that feed algae.[12] • Mischmetal, a pyrophoric alloy used in lighter flints, contains 25% to 45% lanthanum.[2] • Lanthanum oxide additive to tungsten is used in gas tungsten arc welding electrodes, as a substitute for radioactive thorium.[13] [14] • Hydrogen sponge alloys can contain lanthanum. These alloys are capable of storing up to 400 times their own volume of hydrogen gas in a reversible adsorption process. Heat energy is released every time they do so; therefore these alloys have possibilities in energy conservation systems.[2] • Various compounds of lanthanum and other rare-earth elements (oxides, chlorides, etc.) are components of various catalysis, such as petroleum cracking catalysts.[15]
Lanthanum • Lanthanum-barium radiometric dating is used to estimate age of rocks and ores, though the technique has limited popularity.[16] • Lanthanum carbonate was approved as a medication (Fosrenol, Shire Pharmaceuticals) to absorb excess phosphate in cases of end-stage renal failure.[17] • Lanthanum fluoride is used in phosphor lamp coatings. Mixed with europium fluoride, it is also applied in the crystal membrane of fluoride ion-selective electrodes.[5] • Like horseradish peroxidase, lanthanum is used as an electron-dense tracer in molecular biology.[18] • Lanthanum is an intermetallic component of nickel-metal hydride batteries.[19]
Biological role Lanthanum has no known biological role. The element is not absorbed orally, and when injected its elimination is very slow. Lanthanum carbonate was approved as a medication Fosrenol to absorb excess phosphate in cases of end-stage renal failure.[17] While lanthanum has pharmacological effects on several receptors and ion channels, its specificity for the GABA receptor is unique among divalent cations. Lanthanum acts at the same modulatory site on the GABA receptor as zinc- a known negative allosteric modulator. The Lanthanum cation La3+ is a positive allosteric modulator at native and recombinant GABA receptors, increasing open channel time and decreasing desensitization in a subunit configuration dependent manner.[20]
Precautions Lanthanum has a low to moderate level of toxicity, and should be handled with care. In animals, the injection of lanthanum solutions produces glycaemia, low blood pressure, degeneration of the spleen and hepatic alterations.
See also • Lanthanum compounds
Books • The Industrial Chemistry of the Lanthanons, Yttrium, Thorium and Uranium, by R.J. Callow, Pergamon Press 1967 • Extractive Metallurgy of Rare Earths, by C.K. Gupta and N. Krishnamurthy, CRC Press 2005 • Nouveau Traite de Chimie Minerale, Vol. VII. Scandium, Yttrium, Elements des Terres Rares, Actinium, P. Pascal, Editor, Masson & Cie 1959 • Chemistry of the Lanthanons, by R.C. Vickery, Butterworths 1953
8
Lanthanum
9
External links • WebElements.com – Lanthanum
[21]
pnb:مناھتنیل
References [1] Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf). CRC press. 2000. ISBN 0849304814. . [2] C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304814. [3] Tsuchiya, T; Kumashiro, R; Tanigaki, K; Matsunaga, Y; Ishitsuka, Mo; Wakahara, T; Maeda, Y; Takano, Y; Aoyagi, M; Akasaka, T; Liu, Mt; Kato, T; Suenaga, K; Jeong, Js; Iijima, S; Kimura, F; Kimura, T; Nagase, S (Jan 2008). "Nanorods of endohedral metallofullerene derivative.". Journal of the American Chemical Society 130 (2): 450–1. doi: 10.1021/ja710396n (http:/ / dx. doi. org/ 10. 1021/ ja710396n). ISSN 0002-7863 (http:/ / worldcat. org/ issn/ 0002-7863). PMID 18095695. [4] Guan, L; Suenaga, K; Okubo, S; Okazaki, T; Iijima, S (Feb 2008). "Metallic wires of lanthanum atoms inside carbon nanotubes". Journal of the American Chemical Society 130 (7): 2162–3. doi: 10.1021/ja7103069 (http:/ / dx. doi. org/ 10. 1021/ ja7103069). ISSN 0002-7863 (http:/ / worldcat. org/ issn/ 0002-7863). PMID 18225905. edit (http:/ / en. wikipedia. org/ wiki/ Template:cite_doi/ 10. 1021. 2fja7103069) [5] Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds (http:/ / books. google. com/ books?id=Xqj-TTzkvTEC& pg=PA243). McGraw-Hill. pp. 444–446. ISBN 0070494398. . Retrieved 2009-06-06. [6] " Chemical reactions of Lanthanum (https:/ / www. webelements. com/ lanthanum/ chemistry. html)". Webelements. . Retrieved 2009-06-06. [7] " Lighting (http:/ / www. 1911encyclopedia. org/ Lighting)". 11th edition of Encyclopedia Britannica (1911). . Retrieved 2009-06-06. [8] Jason D. Sommerville and Lyon B. King. " Effect of Cathode Position on Hall-Effect Thruster Performance and Cathode Coupling Voltage (http:/ / www. me. mtu. edu/ researchAreas/ isp/ AIAA-2007-5174-907. pdf)". 43rd AIAA/ASME/SAE/ASEE Joint Propulsion Conference & Exhibit, 8 - 11 July 2007, Cincinnati, OH. . Retrieved 2009-06-06. [9] Harrington, James A.. " Infrared Fiber Optics (http:/ / irfibers. rutgers. edu/ pdf_files/ ir_fiber_review. pdf)" (PDF). Rutgers University. . [10] " BrilLanCeTM Scintillators Performance Summary: Scintillation Products Technical Note (http:/ / www. detectors. saint-gobain. com/ Media/ Documents/ S0000000000000001004/ SGC_BrilLance_Scintillators_Performance_Summary. pdf)". . Retrieved 2009-06-06. [11] Kim, K (2003). "The effect of lanthanum on the fabrication of ZrB2–ZrC composites by spark plasma sintering". Materials Characterization 50: 31. doi: 10.1016/S1044-5803(03)00055-X (http:/ / dx. doi. org/ 10. 1016/ S1044-5803(03)00055-X). [12] " Phosphate in swimming pools - the real cause of algae (http:/ / www. articlesbase. com/ home-and-family-articles/ phosphate-in-swimming-pools-the-real-cause-of-algae-860236. html)". . Retrieved 2009-06-06. [13] Howard B. Cary (1995). Arc welding automation (http:/ / books. google. co. jp/ books?id=H3BgQGdTP_0C& ). CRC Press. p. 139. ISBN 0824796454. . [14] Larry Jeffus. (2003). " Types of Tungsten (http:/ / books. google. de/ books?id=zeRiW7en7HAC& pg=RA1-PA750)". Welding : principles and applications. Clifton Park, N.Y.: Thomson/Delmar Learning. p. 350. ISBN 9781401810467. . [15] C. K. Gupta, Nagaiyar Krishnamurthy (2004). Extractive metallurgy of rare earths (http:/ / books. google. co. jp/ books?id=F0Bte_XhzoAC& pg=PA441). CRC Press. p. 441. ISBN 0415333407. . [16] S. Nakai, A. Masuda, B. Lehmann (1988). " La-Ba dating of bastnaesite (http:/ / www. minsocam. org/ ammin/ AM73/ AM73_1111. pdf)". American Mineralogist 7: 1111. . [17] " FDA approves Fosrenol(R) in end-stage renal disease (ESRD) patients (http:/ / www. medicalnewstoday. com/ articles/ 15538. php)". 28 October 2004. . Retrieved 2009-06-06. [18] Chau YP, Lu KS (1995). "Investigation of the blood-ganglion barrier properties in rat sympathetic ganglia by using lanthanum ion and horseradish peroxidase as tracers". Acta Anatomica (Basel) 153 (2): 135–144. PMID 8560966. [19] " Inside the Nickel Metal Hydride Battery (http:/ / www. cobasys. com/ pdf/ tutorial/ inside_nimh_battery_technology. pdf)". . Retrieved 2009-06-06.
Lanthanum [20] Boldyreva, A. A. (2005). "Lanthanum Potentiates GABA-Activated Currents in Rat Pyramidal Neurons of CA1 Hippocampal Field". Bulletin of Experimental Biology and Medicine 140: 403. doi: 10.1007/s10517-005-0503-z (http:/ / dx. doi. org/ 10. 1007/ s10517-005-0503-z). [21] http:/ / www. webelements. com/ webelements/ elements/ text/ La/ index. html
10
Article Sources and Contributors
Article Sources and Contributors Lanthanum Source: http://en.wikipedia.org/w/index.php?oldid=308710611 Contributors: 65.68.87.xxx, Ahoerstemeier, Ale jrb, Alex.muller, Anclation, Archimerged, Arkuat, Beetstra, Belizefan, Benbest, Bobo192, Borislav Dopudja, Bowlhover, Bryan Derksen, CanisRufus, Capricorn42, Carnildo, Chowbok, CommonsDelinker, Conversion script, Cryptic C62, Darrien, David Latapie, DocWatson42, Donarreiskoffer, Download, Dv82matt, Dwmyers, Edgar181, Eleland, Emperorbma, Femto, Fireice, Greatpatton, Grendelkhan, GuelphGryphon98, Hak-kâ-ngìn, Hallpriest9, Hannibal, Helge Skjeveland, Heron, Hugh2414, IanOsgood, Icairns, Ideyal, J.delanoy, Jaan513, Janke, Jaraalbe, Joanjoc, John, Jose77, Kostisl, Kwamikagami, LeaveSleaves, Looxix, Malerin, Materialscientist, Mav, Michael Snow, Minesweeper, Mkweise, Mmm, Moonriddengirl, Mortdefides, Nakon, Nephron, Nergaal, Omicronpersei8, PeepP, Plexust, Poolkris, Pras, PseudoOne, RHB, RTC, Remember, Rjwilmsi, Roberta F., Rominandreu, Salsb, Sam Hocevar, Saperaud, Schneelocke, Sfuerst, Shaddack, Sheitan, Sl, Swedish fusilier, Tagishsimon, Terrace4, Tetracube, Thricecube, Velvetron, Viralmemesis, Vontafeijos, Vsmith, Walkerma, WereSpielChequers, Wrenchelle, Yekrats, Yyy, 141 anonymous edits
Image Sources, Licenses and Contributors image:La-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:La-TableImage.png License: GNU Free Documentation License Contributors: Paddy, Paginazero, Saperaud Image: Lanthan 1-cropflipped.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Lanthan_1-cropflipped.jpg License: unknown Contributors: User:Materialscientist Image:Monazit - Madagaskar.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Monazit_-_Madagaskar.jpg License: unknown Contributors: User:Ra'ike Image:Monazit opening acid.gif Source: http://en.wikipedia.org/w/index.php?title=File:Monazit_opening_acid.gif License: unknown Contributors: User:Hermann Luyken Image:Glowing gas mantle.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Glowing_gas_mantle.jpg License: unknown Contributors: Fourpointsix Image:LaB6HotCathode.jpg Source: http://en.wikipedia.org/w/index.php?title=File:LaB6HotCathode.jpg License: unknown Contributors: Original uploader was Ahecht at en.wikipedia File:Zblan transmit.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Zblan_transmit.jpg License: Public Domain Contributors: NASA
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
11
Cerium
1
Cerium 58
lanthanum ← cerium → praseodymium
↑
Ce ↓
Th Periodic Table - Extended Periodic Table
General Name, Symbol, Number
cerium, Ce, 58
Element category
lanthanides
Group, Period, Block
n/a, 6, f
Appearance
silvery white
Standard atomic weight
140.116(1) g·mol
Electron configuration
[Xe] 4f 5d 6s
−1
1
Electrons per shell
1
2
2, 8, 18, 19, 9, 2 Physical properties
Phase
solid
Density (near r.t.)
6.770 g·cm
Liquid density at m.p.
6.55 g·cm−3
Melting point
1068 K (795 °C, 1463 °F)
Boiling point
3716 K (3443 °C, 6229 °F)
Heat of fusion
5.46 kJ·mol−1
Heat of vaporization
398 kJ·mol−1
Specific heat capacity
(25 °C) 26.94 J·mol−1·K−1
−3
Vapor pressure P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1992
2194
2442
2754
3159
3705
Atomic properties Crystal structure
face-centered cubic
Cerium
2
Oxidation states
4, 3, 2 (mildly basic oxide)
Electronegativity
1.12 (Pauling scale)
Ionization energies (more)
1st: 534.4 kJ·mol−1 2nd: 1050 kJ·mol−1 3rd: 1949 kJ·mol−1
Atomic radius
181.8 pm
Covalent radius
204±9 pm Miscellaneous [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(r.t.) (β, poly) 828 nΩ·m
Thermal conductivity
(300 K) 11.3 W·m
Thermal expansion
(r.t.) (γ, poly) 6.3 µm/(m·K)
Speed of sound (thin rod)
(20 °C) 2100 m/s
Young's modulus
(γ form) 33.6 GPa
Shear modulus
(γ form) 13.5 GPa
Bulk modulus
(γ form) 21.5 GPa
Poisson ratio
(γ form) 0.24
Mohs hardness
2.5
Vickers hardness
270 MPa
Brinell hardness
412 MPa
CAS registry number
−1
·K
−1
7440-45-1 Most-stable isotopes
Main article: Isotopes of cerium iso 134
Ce
NA syn
half-life 3.16 days
136
0.185%
136
138
0.251%
138
Ce Ce
139
Ce
140
Ce
141
Ce
142
Ce
144
Ce
syn 88.450% syn 11.114% syn
DM ε
DE (MeV)
DP
0.500
134
0.278
139
La
Ce is stable with 78 neutron Ce is stable with 80 neutron
137.640 days
ε
La
140
Ce is stable with 82 neutron
32.501 days
β−
0.581
141
> 5×1016 years
β−β−
1.417
142
284.893 days
β−
0.319
144
References
Pr Nd Pr
Cerium
3
Cerium (pronounced /ˈsɪəriəm/) is a chemical element with the symbol Ce and atomic number 58. It is a soft, silvery, ductile metal which easily oxidizes in air. Cerium was named after the dwarf planet Ceres. Cerium is the most abundant of the rare earth elements, making up about 0.0046% of the Earth's crust by weight. It is found in a number of minerals, the most important being monazite and bastnasite. Commercial applications of cerium are numerous. They include catalysts, additives to fuel to reduce emissions and to glass and enamels to change their color. Cerium oxide is an important component of glass polishing powders and phosphors used in screens and fluorescent lamps.
Characteristics Physical Cerium is a silvery metal, belonging to the lanthanoid group. It resembles iron in color and luster, but is soft, and both malleable and ductile. Cerium has the longest liquid range of any non-radioactive element: 2648 C° (795 °C to 3443 °C) or 4766 F° (1463 °F to 6229 °F). (Thorium has a longer liquid range, but is radioactive) Cerium is especially interesting because of its variable electronic structure. The energy of the inner 4f level is nearly the same as that of the outer or valence electrons, and only small energy is required to change the relative occupancy of these electronic levels. This gives rise to dual valency states. For example, a volume change of about 10% occurs when cerium is subjected to high pressures or low temperatures. It appears that the Phase diagram of cerium valence changes from about 3 to 4 when it is cooled or compressed. The low temperature behavior of cerium is complex. Four allotropic modifications are thought to exist: cerium at room temperature and at atmospheric pressure is known as γ cerium. Upon cooling to –16°C, γ cerium changes to ß cerium. The remaining γ cerium starts to change to α cerium when cooled to –172°C, and the transformation is complete at –269 °C. α Cerium has a density of 8.16; δ cerium exists above 726 °C. At atmospheric pressure, liquid cerium is more dense than its solid form at the melting point.[2] [3] [4]
Cerium
4
Chemical Cerium metal tarnishes slowly in air and burns readily at 150 °C to form cerium(IV) oxide: Ce + O2 → CeO2 Cerium is quite electropositive and reacts slowly with cold water and quite quickly with hot water to form cerium hydroxide: 2 Ce (s) + 6 H2O (g) → 2 Ce(OH)3 (aq) + 3 H2 (g) Cerium metal reacts with all the halogens: 2 Ce (s) + 3 F2 (g) → 2 CeF3 (s) [white] 2 Ce (s) + 3 Cl2 (g) → 2 CeCl3 (s) [white] 2 Ce (s) + 3 Br2 (g) → 2 CeBr3 (s) [white] 2 Ce (s) + 3 I2 (g) → 2 CeI3 (s) [yellow] Cerium dissolves readily in dilute sulfuric acid to form solutions containing the colorless Ce(III) ions, which exist as a [Ce(OH2)9]3+ complexes:[5] 2 Ce (s) + 3 H2SO4 (aq) → 2 Ce3+ (aq) + 3 SO2−4 (aq) + 3 H2 (g)
Compounds Cerium(IV) (ceric) salts are orange red or yellowish, whereas cerium(III) (cerous) salts are usually white or colorless. Both oxidation states absorb ultraviolet light strongly. Cerium(III) can be used to make glasses that are colorless, yet absorb ultraviolet light almost completely. Cerium can be readily detected in rare earth mixtures by a very sensitive qualitative test: addition of ammonia and hydrogen peroxide to an aqueous solution of lanthanides produces a characteristic dark brown color if cerium is present.
Cerium(IV) sulfate
Cerium exhibits three oxidation states, +2, +3 and +4. The +2 state is rare and is observed in CeH2, CeI2 and CeS.[4] The most common compound of cerium is cerium(IV) oxide (CeO2), which is used as "jeweller's rouge" as well as in the walls of some self-cleaning ovens. Two common oxidizing agents used in titrations are ammonium cerium(IV) sulfate (ceric ammonium sulfate, (NH4)2Ce(SO4)3) and ammonium cerium(IV) nitrate (ceric ammonium nitrate or CAN, (NH4)2Ce(NO3)6). Cerium also forms a chloride, CeCl3 or cerium(III) chloride, used to facilitate reactions at carbonyl groups in organic chemistry. Other compounds include cerium(III) carbonate (Ce2(CO3)3), cerium(III) fluoride (CeF3), cerium(III) oxide (Ce2O3), as well as cerium(IV) sulfate (ceric sulfate, Ce(SO4)2) and cerium(III) triflate (Ce(OSO2CF3)3). The two oxidation states of cerium differ enormously in basicity: cerium(III) is a strong base, comparable to the other trivalent lanthanides, but cerium(IV) is weak. This difference has always allowed cerium to be by far the most readily isolated and purified of all the lanthanides, otherwise a notoriously difficult group of elements to separate. A wide range of procedures have been devised over the years to exploit the difference. Among the better ones:
Cerium 1. Leaching the mixed hydroxides with dilute nitric acid: the trivalent lanthanides dissolve in cerium-free condition, and tetravalent cerium remains in the insoluble residue as a concentrate to be further purified by other means. A variation on this uses hydrochloric acid and the calcined oxides from bastnasite, but the separation is less sharp. 2. Precipitating cerium from a nitrate or chloride solution using potassium permanganate and sodium carbonate in a 1:4 molar ratio. 3. Boiling rare-earth nitrate solutions with potassium bromate and marble chips. Formerly used commercially was a method whereby a solution of cerium(IV) in nitric acid would be added to dilute sulfuric acid. This caused cerium(IV) to largely precipitate as a basic salt, leaving trivalent lanthanide in solution. However, the finely divided precipitate was difficult to filter from the highly corrosive medium. Using the classical methods of rare-earth separation, there was a considerable advantage to a strategy of removing cerium from the mixture at the beginning. Cerium typically comprised 45% of the cerite or monazite rare earths, and removing it early greatly reduced the bulk of what needed to be further processed (or the cost of reagents to be associated with such processing). However, not all cerium purification methods relied on basicity. Ceric ammonium nitrate [ammonium hexanitratocerate(IV)] crystallization from nitric acid was one purification method. Cerium(IV) nitrate (hexanitratoceric acid) was more readily extractable into certain solvents (e.g. tri-n-butyl phosphate) than the trivalent lanthanides. However, modern practice in China seems to be to do purification of cerium by counter-current solvent extraction, in its trivalent form, just like the other lanthanides. Cerium(IV) is a strong oxidant under acidic conditions, but stable under alkaline conditions, when it is cerium(III) that becomes a strong reductant, easily oxidized by atmospheric oxygen (O2). This ease of oxidation under alkaline conditions leads to the occasional geochemical parting of the ways between cerium and the trivalent light lanthanides under supergene weathering conditions, leading variously to the "negative cerium anomaly" or to the formation of the mineral cerianite. Air-oxidation of alkaline cerium(III) is the most economical way to get to cerium(IV), which can then be handled in acid solution.
Isotopes Naturally-occurring cerium is composed of 4 stable isotopes; 136Ce, 138Ce, 140Ce, and 142Ce with 140Ce being the most abundant (88.48% natural abundance). 136Ce and 142Ce are predicted to be double beta active but no signs of activity were ever observed (for 142Ce, the lower limit on half-life is 5×1016 years). 26 radioisotopes have been characterized with the most long-lived being 144Ce with a half-life of 284.893 days, 139Ce with a half-life of 137.640 days, and 141Ce with a half-life of 32.501 days. All of the remaining radioactive isotopes have half-lives that are less than 4 days and the majority of these have half-lives that are less than 10 minutes. This element also has 2 meta states. The known isotopes of cerium range in atomic weight from 123 u (123Ce) to 152 u (152Ce). Cerium 144 is a high-yield product of nuclear fission; the ORNL Fission Product Pilot Plant separated substantial quantities of cerium-144 from reactor waste, and it was used in the Aircraft Nuclear Propulsion and SNAP programs.
5
Cerium
History Cerium was discovered in Bastnäs in Sweden by Jöns Jakob Berzelius and Wilhelm Hisinger, and independently in Germany by Martin Heinrich Klaproth, both in 1803. Cerium was so named by Berzelius after the dwarf planet Ceres, discovered two years earlier (1801). As originally isolated, cerium was in the form of its oxide, and was named ceria, a term that is still used. The metal itself was too electropositive to be isolated by then-current smelting technology, a characteristic of rare earth metals in general. However, the development of electrochemistry by Humphry Davy was only five years into the future, and then the earths were soon to yield the metals they contained. Ceria, as isolated in 1803, contained all of the lanthanides present in the cerite ore from Bastnäs, Sweden, and thus only contained about 45% of what is now known to be pure ceria. It was not until Carl Gustaf Mosander succeeded in removing lanthana and "didymia" in the late 1830s, that ceria was obtained pure. As a historical aside: Wilhelm Hisinger was a wealthy mine owner and amateur scientist, and sponsor of Berzelius. He owned or controlled the mine at Bastnäs, and had been trying for years to find out the composition of the abundant heavy gangue rock (the "Tungstein of Bastnäs"), now known as cerite, that he had in his mine. Mosander and his family lived for many years in the same house as Berzelius, and the former was undoubtedly persuaded by the latter to investigate ceria further. When the rare earths were first discovered, since they were strong bases like the oxides of calcium or magnesium, they were thought to be divalent. Thus, "ceric" cerium was thought to be trivalent, and the oxidation state ratio was therefore thought to be 1.5. Berzelius was extremely annoyed to keep on getting the ratio 1.33. He was after all one of the finest analytical chemists in Europe. But he was a better analyst than he thought, since 1.33 was the correct answer! In the late 1950s, the Lindsay Chemical Division of American Potash and Chemical Corporation of West Chicago, Illinois, then the largest producer of rare earths in the world, was offering cerium compounds in two purity ranges, "commercial" at 94-97% purity, and "purified", at a reported 99.9+% purity. In their October 1, 1958 pricelist, one-pound quantities of the oxides were priced at $3.30 or $8.10 respectively for the two purities; the per-pound price for 50-pound quantities were respectively $1.95 or $4.95 for the two grades. Cerium salts were proportionately cheaper, reflecting their lower net content of oxide.
6
Cerium
7
Occurrence Cerium is the most abundant of the rare earth elements, making up about 0.0046% of the Earth's crust by weight. It is found in a number of minerals including allanite (also known as orthite)—(Ca, Ce, La, Y)2(Al, Fe)3(SiO4)3(OH), monazite (Ce, La, Th, Nd, Y)PO4, bastnasite (Ce, La, Y)CO3F, hydroxylbastnasite (Ce, La, Nd)CO3(OH, F), rhabdophane (Ce, La, Nd)PO4-H2O, zircon (ZrSiO4), and synchysite Ca(Ce, La, Nd, Y)(CO3)2F. Monazite and bastnasite are presently the two most important sources of cerium. Large deposits of monazite, allanite, and bastnasite will supply cerium, thorium, and other rare-earth metals for many years to come.[3]
Allanite
Production The mineral mixtures are crushed, ground and treated with hot concentrated sulfuric acid to produce water-soluble sulfates of rare earths. The acidic filtrates are partially neutralized with sodium hydroxide to pH 3-4. Thorium precipitates out of solution as hydroxide and is removed. After that the solution is treated with ammonium oxalate to convert rare earths in to their insoluble oxalates. The oxalates are converted to oxides by annealing. The oxides are dissolved in nitric acid that excludes one of the main components, cerium, whose salts is insoluble in HNO3. Metallic cerium is prepared by metallothermic reduction techniques, such as by reducing cerium fluoride or chloride with calcium, or by electrolysis of molten cerous chloride or other cerous halides. The metallothermic technique is used to produce high-purity cerium. [4]
Applications A major technological application for Cerium(III) oxide is a catalytic converter for the reduction of CO emissions in the exhaust gases from motor vehicles. In particular, cerium oxide is added into Diesel fuels. Another important use of the cerium oxide is a hydrocarbon catalyst in self cleaning ovens, incorporated into oven walls and as a petroleum cracking catalyst in petroleum refining. [6] Cerium(IV) oxide is considered one of the most efficient agents for precision polishing of optical components. Cerium compounds are also used in the manufacture of glass, both as a component and as a decolorizer. For example, cerium(IV) oxide in combination with titanium(IV) oxide gives a golden yellow color to glass; it also allows for selective absorption of ultraviolet light in glass. Cerium oxide has high refractive index and is added to enamel to make it more opaque. [6] Cerium(IV) oxide is used in incandescent gas mantles, such as the Welsbach mantle, where it was combined with thorium, lanthanum, magnesium or yttrium oxides. Doped with other rare earth oxides, it has been investigated as a solid electrolyte in intermediate temperature solid oxide fuel cells: The cerium(IV) oxide-cerium(III) oxide cycle or CeO2/Ce2O3 cycle is a two step thermochemical process based on cerium(IV) oxide and
Cerium cerium(III) oxide for hydrogen production[7] . The photostability of pigments can be enhanced by addition of cerium. It provides pigments with light fastness and prevents clear polymers from darkening in sunlight. Television glass plates are subject to electron bombardment, which tends to darken them by creation of F-center color centers. This effect is suppressed by addition of cerium oxide. Cerium is also an essential component of phosphors used in TV screens and fluorescent lamps.[6] A traditional use of cerium was in the pyrophoric mischmetal alloy used for light flints. Because of the high affinity of cerium to sulfur and oxygen, it is used in various aluminium alloys, and iron alloys. In steels, cerium degasifies and can help reduce sulfides and oxides, and it is a precipitation hardening agent in stainless steel. Adding cerium to cast irons opposes graphitization and produces a malleable iron. Addition of 3-4% of cerium to magnesium alloys, along with 0.2 to 0.6% zirconium, helps refine the grain and give sound casting of complex shapes. It also adds heat resistance to magnesium castings.[6] Cerium alloys are used in permanent magnets and in tungsten electrodes for gas tungsten arc welding. Cerium is used in carbon-arc lighting, especially in the motion picture industry. Cerium oxalate is an anti-emetic drug. Cerium(IV) sulfate is used extensively as a volumetric oxidizing agent in quantitative analysis. Ceric ammonium nitrate is a useful one-electron oxidant in organic chemistry, used to oxidatively etch electronic components, and as a primary standard for quantitative analysis.[3] [8]
Precautions Cerium, like all rare-earth metals, is of low to moderate toxicity. Cerium is a strong reducing agent and ignites spontaneously in air at 65 to 80 °C. Fumes from cerium fires are toxic. Water should not be used to stop cerium fires, as cerium reacts with water to produce hydrogen gas. Workers exposed to cerium have experienced itching, sensitivity to heat, and skin lesions. Animals injected with large doses of cerium have died due to cardiovascular collapse. Cerium(IV) oxide is a powerful oxidizing agent at high temperatures and will react with combustible organic materials. While cerium is not radioactive, the impure commercial grade may contain traces of thorium, which is radioactive. Cerium serves no known biological function.[6]
External links • WebElements.com – Cerium [9] • It's Elemental – The Element Cerium [10] • Cerium Properties and Applications [11] pnb:میریس
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] Stassis, C. (1979). "Lattice and spin dynamics of γ-Ce". Physical Review B 19: 5746. doi: 10.1103/PhysRevB.19.5746 (http:/ / dx. doi. org/ 10. 1103/ PhysRevB. 19. 5746). [3] C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304814. [4] Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds (http:/ / books. google. com/ books?id=Xqj-TTzkvTEC& pg=PA200). McGraw-Hill. pp. 199–200. ISBN 0070494398. . Retrieved 2009-06-06.
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Cerium [5] " Chemical reactions of Cerium (https:/ / www. webelements. com/ cerium/ chemistry. html)". Webelements. . Retrieved 2009-06-06. [6] Alessandro (2002). Catalysis by ceria and related materials (http:/ / books. google. com/ books?id=X2z9WdN3-WgC). Imperial College Press. pp. 6–11. ISBN 1860942997. . [7] " Hydrogen production from solar thermochemical water splitting cycles (http:/ / www. solarpaces. org/ Tasks/ Task2/ HPST. HTM)". . Retrieved 2009-06-06. [8] C. K. Gupta, Nagaiyar Krishnamurthy (2004). Extractive metallurgy of rare earths (http:/ / books. google. co. jp/ books?id=F0Bte_XhzoAC& pg=PA30). CRC Press. p. 30. ISBN 0415333407. . [9] http:/ / www. webelements. com/ webelements/ elements/ text/ Ce/ index. html [10] http:/ / education. jlab. org/ itselemental/ ele058. html [11] http:/ / www. azom. com/ details. asp?ArticleID=592
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Article Sources and Contributors
10
Article Sources and Contributors Cerium Source: http://en.wikipedia.org/w/index.php?oldid=308637771 Contributors: 65.68.87.xxx, Acroterion, Ahoerstemeier, Alansohn, Andycjp, Archimerged, Arkuat, AubreyEllenShomo, Baccyak4H, Beavertank, Beetstra, Benbest, Bensaccount, BillFlis, BlueEarth, Bobblewik, Borislav Dopudja, Bryan Derksen, CanisRufus, Carlosguitar, Carnildo, CommonsDelinker, Conversion script, Cuenca, Darrien, David Latapie, DerHexer, Donarreiskoffer, Doulos Christos, Dryguy, Duk, Dwmyers, Edgar181, Emperorbma, Epbr123, Eranb, Femto, Fivemack, Gaz, Gene Nygaard, Geoking66, Giftlite, Greatpatton, Grendelkhan, Grimlock, Hallpriest9, Icairns, Ideyal, Ilikepie2221, Ivan05, Jaganath, Janke, Jaraalbe, Jll, Joanjoc, John, Jons63, Jose77, Kumorifox, Kurykh, Kwamikagami, Lightmouse, Lkitrossky, Looxix, Manscher, Materialscientist, Mav, Mdf, Megan1967, Michael Snow, Minesweeper, Mion, Mortdefides, NawlinWiki, Nergaal, Nihiltres, NuclearWarfare, Ozzieboy, Paul D. Anderson, Pavel Vozenilek, Pdn, PeepP, Plantsurfer, Polyparadigm, Poolkris, Pusher, RTC, Red1530, Remember, Richard B, Rifleman 82, Roberta F., Rocastelo, Romanm, Ronbo76, Rror, SDC, Saaga, Saperaud, Schneelocke, Sengkang, Sfuerst, Shinkolobwe, Skizzik, Sl, Snoyes, SpK, Spangineer, StealthFox, Stifynsemons, Tagishsimon, Tetracube, Thricecube, Tiddly Tom, Titoxd, Trevor MacInnis, Tullywinters, V1adis1av, V8rik, Velvetron, Vsmith, Vuo, Walkerma, Warut, Watch37264, Wizard191, Xcomradex, Yekrats, Yyy, ZX81, 158 anonymous edits
Image Sources, Licenses and Contributors image:Ce-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Ce-TableImage.png License: GNU Free Documentation License Contributors: Conscious, Georg Slickers, Paddy, Paginazero, Saperaud Image: CE2k2g-crop.jpg Source: http://en.wikipedia.org/w/index.php?title=File:CE2k2g-crop.jpg License: unknown Contributors: User:Materialscientist Image:Cerium phase diagram.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Cerium_phase_diagram.jpg License: GNU Free Documentation License Contributors: User:Grimlock Image:Cer(IV)-sulfat.JPG Source: http://en.wikipedia.org/w/index.php?title=File:Cer(IV)-sulfat.JPG License: GNU Free Documentation License Contributors: BXXXD
Image:Allanite w-rock Basic calcium iron Rare Earth aluminum silicate near 29 Palms San Bernardino County California 2204.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Allanite_w-rock_Basic_calcium_iron_Rare_Earth_aluminum_silicate_near_29_Palms_San_Bernardino_County_California_2204.jpg License: Public Domain Contributors: Dave Dyet http://www.shutterstone.com http://www.dyet.com
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
Praseodymium
1
Praseodymium 59
cerium ← praseodymium → neodymium
↑
Pr ↓
Pa Periodic Table - Extended Periodic Table
General Name, Symbol, Number
praseodymium, Pr, 59
Element category
lanthanides
Group, Period, Block
n/a, 6, f
Appearance
grayish white
Standard atomic weight
140.90765(2) g·mol
Electron configuration
[Xe] 4f 6s
−1
3
Electrons per shell
2
2, 8, 18, 21, 8, 2 Physical properties
Phase
solid
Density (near r.t.)
6.77 g·cm−3
Liquid density at m.p.
6.50 g·cm−3
Melting point
1208 K (935 °C, 1715 °F)
Boiling point
3793 K (3520 °C, 6368 °F)
Heat of fusion
6.89 kJ·mol−1
Heat of vaporization
331 kJ·mol−1
Specific heat capacity
(25 °C) 27.20 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1771
1973
(2227)
(2571)
(3054)
(3779)
Atomic properties
Praseodymium
2
Crystal structure
hexagonal
Oxidation states
4, 3, 2 (mildly basic oxide)
Electronegativity
1.13 (Pauling scale)
Ionization energies (more)
1st: 527 kJ·mol−1 2nd: 1020 kJ·mol−1 3rd: 2086 kJ·mol−1
Atomic radius
182 pm
Covalent radius
203±7 pm Miscellaneous [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(r.t.) (α, poly) 0.700 µΩ·m
Thermal conductivity
(300 K) 12.5 W·m
Thermal expansion
(r.t.) (α, poly) 6.7 µm/(m·K)
Speed of sound (thin rod)
(20 °C) 2280 m/s
Young's modulus
(α form) 37.3 GPa
Shear modulus
(α form) 14.8 GPa
Bulk modulus
(α form) 28.8 GPa
Poisson ratio
−1
·K
−1
(α form) 0.281
Vickers hardness
400 MPa
Brinell hardness
481 MPa
CAS registry number
7440-10-0 Most-stable isotopes
Main article: Isotopes of praseodymium iso 141
Pr
142
Pr
143
Pr
NA 100% syn
syn
half-life
DM
DE (MeV)
DP
141
Pr is stable with 82 neutron
19.12 h
13.57 d
β−
2.162
142
ε
0.745
142
β−
0.934
143
Nd Ce Nd
References
Praseodymium (pronounced /ˌpreɪzi.ɵˈdɪmiəm/ or /ˌpreɪsioʊˈdɪmiəm/) is a chemical element that has the symbol Pr and atomic number 59.
Praseodymium
Characteristics Physical Praseodymium is a soft, silvery, malleable and ductile metal in the lanthanoid group. It is somewhat more resistant to corrosion in air than europium, lanthanum, cerium, or neodymium, but it does develop a green oxide coating that spalls off when exposed to air, exposing more metal to oxidation. For this reason, praseodymium should be stored under a light mineral oil or sealed in glass. Contrary to other rare-earth metals, which show antiferromagnetic or/and ferromagnetic ordering at low temperatures, Pr is paramagnetic at any temperatures above 1 K.[1]
Chemical Praseodymium metal tarnishes slowly in air and burns readily at 150 °C to form praseodymium(III,IV) oxide: 12 Pr + 11 O2 → 2 Pr6O11 Praseodymium is quite electropositive and reacts slowly with cold water and quite quickly with hot water to form praseodymium hydroxide: 2 Pr (s) + 6 H2O (g) → 2 Pr(OH)3 (aq) + 3 H2 (g) Praseodymium metal reacts with all the halogens: 2 Pr (s) + 3 F2 (g) → 2 PrF3 (s) [green] 2 Pr (s) + 3 Cl2 (g) → 2 PrCl3 (s) [green] 2 Pr (s) + 3 Br2 (g) → 2 PrBr3 (s) [green] 2 Pr (s) + 3 I2 (g) → 2 PrI3 (s) Praseodymium dissolves readily in dilute sulphuric acid to form solutions containing green Pr(III) ions, which exist as a [Pr(OH2)9]3+ complexes:[2] 2 Pr (s) + 3 H2SO4 (aq) → 2 Pr3+(aq) + 3 SO2−4 (aq) + 3 H2 (g)
Compounds In its compounds, praseodymium occurs in oxidation states +2, +3 and/or +4. Praseodymium(IV) is a strong oxidant, instantly oxidizing water to elemental oxygen (O2), or hydrochloric acid to elemental chlorine. Thus, in aqueous solution, only the +3 oxidation state is encountered. Praseodymium(III) salts are yellow-green and, in solution, present a fairly simple absorption spectrum in the visible region, with a band in the yellow-orange at 589-590 nm (which coincides with the sodium emission doublet), and three bands in the blue/violet region, at 444, 468, and 482 nm, approximately. These positions vary slightly with the counter-ion. Praseodymium oxide, as obtained by the ignition of salts such as the oxalate or carbonate in air, is essentially black in color (with a hint of brown or green) and contains +3 and +4 praseodymium in a somewhat variable ratio, depending upon the conditions of formation. Its formula is conventionally rendered as Pr6O11. Other praseodymium compounds include: • Fluorides: PrF2, PrF3, PrF4 • Chlorides: PrCl3 • Bromides: PrBr3, Pr2Br5 • Iodides: PrI2, PrI3, Pr2I5
3
Praseodymium • • • • • •
Oxides: PrO2, Pr2O3, Pr6O11 Sulfides: PrS, Pr2S3 Sulfates: Pr2(SO4)3 Selenides: PrSe Tellurides: PrTe, Pr2Te3 Nitrides: PrN
Isotopes Naturally occurring praseodymium is composed of one stable isotope, 141Pr. Thirty-eight radioisotopes have been characterized with the most stable being 143Pr with a half-life of 13.57 days and 142Pr with a half-life of 19.12 hours. All of the remaining radioactive isotopes have half-lives that are less than 5.985 hours and the majority of these have half-lives that are less than 33 seconds. This element also has six meta states with the most stable being 138mPr (t½ 2.12 hours), 142mPr (t½ 14.6 minutes) and 134mPr (t½ 11 minutes). The isotopes of praseodymium range in atomic weight from 120.955 u (121Pr) to 158.955 u (159Pr). The primary decay mode before the stable isotope, 141Pr, is electron capture and the primary mode after is beta minus decay. The primary decay products before 141Pr are element 58 (cerium) isotopes and the primary products after are element 60 (neodymium) isotopes.
History The name praseodymium comes from the Greek prasios, meaning green, and didymos, twin. Praseodymium is frequently misspelled as praseodynium. In 1841, Mosander extracted the rare earth didymium from lanthana. In 1874, Per Teodor Cleve concluded that didymium was in fact two elements, and in 1879, Lecoq de Boisbaudran isolated a new earth, samarium, from didymium obtained from the mineral samarskite. In 1885, the Austrian chemist baron Carl Auer von Welsbach separated didymium into two elements, praseodymium and neodymium, which gave salts of different colors. Leo Moser (son of Ludwig Moser, founder of the Moser Glassworks in what is now Karlovy Vary, Bohemia, in the Czech Republic, not to be confused with Leo Moser, a mathematician) investigated the use of praseodymium in glass coloration in the late 1920s. The result was a yellow-green glass given the name "Prasemit". However, a similar color could be achieved with colorants costing only a minute fraction of what praseodymium cost in the late 1920s, such that the color was not popular, few pieces were made, and examples are now extremely rare. Moser also blended praseodymium with neodymium to produce "Heliolite" glass ("Heliolit" in German), which was more widely accepted. The first enduring commercial use of purified praseodymium, which continues today, is in the form of a yellow-orange stain for ceramics, "Praseodymium Yellow", which is a solid-solution of praseodymium in the zirconium silicate (zircon) lattice. This stain has no hint of green in it. By contrast, at sufficiently high loadings, praseodymium glass is distinctly green, rather than pure yellow. Using classical separation methods, praseodymium was always difficult to purify. Much less abundant than the lanthanum and neodymium from which it was being separated (cerium having long since been removed by redox chemistry), praseodymium ended up being dispersed among a large number of fractions, and the resulting yields of purified material
4
Praseodymium
5
were low. Praseodymium has historically been a rare earth whose supply has exceeded demand. This has occasionally led to its being offered more cheaply than the far more abundant neodymium. Unwanted as such, much praseodymium has been marketed as a mixture with lanthanum and cerium, or "LCP" for the first letters of each of the constituents, for use in replacing the traditional lanthanide mixtures that were inexpensively made from monazite or bastnaesite. LCP is what remains of such mixtures, after the desirable neodymium, and all the heavier, rarer and more valuable lanthanides have been removed, by solvent extraction. However, as technology progresses, praseodymium has been found possible to incorporate into neodymium-iron-boron magnets, thereby extending the supply of the much in demand neodymium. So LC is starting to replace LCP as a result.
Occurrence Praseodymium is available in small quantities in Earth’s crust (9.5 ppm). It is found in the rare earth minerals monazite and bastnasite, typically comprising about 5% of the lanthanides contained therein, and can be recovered from bastnasite or monazite by an ion exchange process, or by counter-current solvent extraction. Misch metal, used in making cigarette lighters, contains about 5% praseodymium metal.[3] Monazite
Applications Uses of praseodymium: • As an alloying agent with magnesium to create high-strength metals that are used in aircraft engines.[4] • Praseodymium forms the core of carbon arc lights which are used in the motion picture industry for studio lighting and projector lights. • Praseodymium compounds give glasses and enamels a yellow color.[5] • Praseodymium is used to color cubic zirconia yellow-green, to simulate mineral peridot. • Praseodymium is a component of didymium glass, which is used to make certain types of welder's and glass blower's goggles.[5] • Silicate glass doped with praseodymium ions has been used to slow a light pulse down to a few hundred meters per second.[6] • Praseodymium alloyed with nickel (PrNi5) has such a strong magnetocaloric effect that it has allowed scientists to approach within one thousandth of a degree of absolute zero.[7] • Doping praseodymium in fluoride glass allows it to be used as a single mode fiber optical amplifier.[8] • Praseodymium oxide in solid solution with ceria, or with ceria-zirconia, have been used as oxidation catalysts.[9] • Modern ferrocerium firesteel products, commonly referred to as "flint," used in lighters, torch strikers, "flint and steel" fire starters, etc., contain around 4% praseodymium.[5]
Praseodymium
6
Precautions Like all rare earths, praseodymium is of low to moderate toxicity. Praseodymium has no known biological role.
Books • R.J. Callow, "The Industrial Chemistry of the Lanthanons, Yttrium, Thorium and Uranium", Pergamon Press, 1967.
External links • WebElements.com – Praseodymium [10] • It's Elemental – The Element Praseodymium
[11]
pnb:میموڈیزیرپ
References [1] M. Jackson "Magnetism of Rare Earth" The IRM quaterly col. 10, No. 3, p. 1, 2000 (http:/ / www. irm. umn. edu/ quarterly/ irmq10-3. pdf) [2] " Chemical reactions of Praseodymium (https:/ / www. webelements. com/ praseodymium/ chemistry. html)". Webelements. . Retrieved 2009-06-06. [3] Gschneidner, K.A., and Eyring, L., Handbook on the Physics and Chemistry of Rare Earths, North Holland Publishing Co., Amsterdam, 1978. [4] L. L. Rokhlin (2003). Magnesium alloys containing rare earth metals: structure and properties. CRC Press. ISBN 0415284147. [5] C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304814. [6] " ANU team stops light in quantum leap (http:/ / info. anu. edu. au/ ovc/ Media/ Media_Releases/ 2005/ August/ 290805_stop_light)". . Retrieved 18 May 2009. [7] Emsley, John (2001). Nature's building blocks. Oxford University Press. pp. 342. ISBN 0-1985-0341-5. [8] Jha, A; Naftaly, M; Jordery, S; Samson, B N; Taylor, E R; Hewak, D; Payne, D N; Poulain, M; Zhang, G (1995). Pure and Applied Optics: Journal of the European Optical Society Part A 4: 417. doi: 10.1088/0963-9659/4/4/019 (http:/ / dx. doi. org/ 10. 1088/ 0963-9659/ 4/ 4/ 019). [9] Borchert, Y.; Sonstrom, P.; Wilhelm, M.; Borchert, H.; Baumer, M. (2008). "Nanostructured Praseodymium Oxide: Preparation, Structure, and Catalytic Properties". Journal of Physical Chemistry C 112: 3054. doi: 10.1021/jp0768524 (http:/ / dx. doi. org/ 10. 1021/ jp0768524). [10] http:/ / www. webelements. com/ webelements/ elements/ text/ Pr/ index. html [11] http:/ / education. jlab. org/ itselemental/ ele059. html
Article Sources and Contributors
Article Sources and Contributors Praseodymium Source: http://en.wikipedia.org/w/index.php?oldid=308867016 Contributors: 234n5 &, 2over0, 65.68.87.xxx, A new name 2008, Aeusoes1, Ahoerstemeier, AlimanRuna, Anthony Appleyard, Archimerged, Arkuat, B07, Baccyak4H, Beano, Beetstra, Benbest, Biblbroks, Borislav Dopudja, Brentdax, Bryan Derksen, Cacycle, Carnildo, Chem-awb, Conversion script, Corpx, DMacks, DaGizza, Darrien, David Latapie, Deli nk, Donarreiskoffer, Dwmyers, Edgar181, Emperorbma, Encyclopedia77, Enotdetcelfer, Fanghong, Femto, Flyguy649, GeeJo, Giftlite, Grendelkhan, Hak-kâ-ngìn, Hazel77, Heron, Icairns, Ideyal, J.delanoy, Jaan513, Jaraalbe, Joanjoc, John, Jose77, Kimse, Kukini, Kumorifox, Kurykh, Kwamikagami, LeaveSleaves, Looxix, Lucinos, Marc Venot, Markjoseph125, Materialscientist, Mav, Mdf, Meelar, Metallica823, Minesweeper, Mmm, Mortdefides, MrFish, Neodymiferous, Nergaal, Nescio, Nihiltres, Oxymoron83, Pacobob, Peak, Plexust, Poolkris, Pras, RTC, Remember, Reyk, Rholton, Roberta F., Romanm, Saperaud, Schneelocke, Sengkang, Shaddack, Sl, Smalljim, Snigbrook, Spidey71, Stifynsemons, Stone, Svante, Swedish fusilier, Tagishsimon, Tartarus, Tetracube, Tohd8BohaithuGh1, Ubergeekguy, VASANTH S.N., Vsmith, Walkerma, Warut, Watch37264, Ww2censor, Yekrats, Yyy, 139 anonymous edits
Image Sources, Licenses and Contributors image:Pr-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Pr-TableImage.png License: GNU Free Documentation License Contributors: user:schneelocke Image: Praseodym 1-cropflipped.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Praseodym_1-cropflipped.jpg License: unknown Contributors: User:Materialscientist File:Monazite - Rostadheia, Iveland, Norvegia 01.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Monazite_-_Rostadheia,_Iveland,_Norvegia_01.jpg License: unknown Contributors: User:Aangelo
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7
Neodymium
1
Neodymium 60
praseodymium ← neodymium → promethium
↑
Nd ↓
U Periodic Table - Extended Periodic Table
General Name, Symbol, Number
neodymium, Nd, 60
Element category
lanthanides
Group, Period, Block
n/a, 6, f
Appearance
silvery white, yellowish tinge
Standard atomic weight
144.242(3) g·mol
Electron configuration
[Xe] 4f 6s
−1
4
Electrons per shell
2
2, 8, 18, 22, 8, 2 Physical properties
Phase
solid
Density (near r.t.)
7.01 g·cm−3
Liquid density at m.p.
6.89 g·cm−3
Melting point
1297 K (1024 °C, 1875 °F)
Boiling point
3347 K (3074 °C, 5565 °F)
Heat of fusion
7.14 kJ·mol−1
Heat of vaporization
289 kJ·mol−1
Specific heat capacity
(25 °C) 27.45 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1595
1774
1998
(2296)
(2715)
(3336)
Atomic properties
Neodymium
2
Crystal structure
hexagonal
Oxidation states
3, 2 (mildly basic oxide)
Electronegativity
1.14 (Pauling scale)
Ionization energies (more)
1st: 533.1 kJ·mol−1 2nd: 1040 kJ·mol−1 3rd: 2130 kJ·mol−1
Atomic radius
181 pm
Covalent radius
201±6 pm Miscellaneous
Magnetic ordering
paramagnetic, antiferromagnetic below 20K
Electrical resistivity
(r.t.) (α, poly) 643 nΩ·m
Thermal conductivity
(300 K) 16.5 W·m
Thermal expansion
(r.t.) (α, poly) 9.6 µm/(m·K)
Speed of sound (thin rod)
(20 °C) 2330 m/s
Young's modulus
(α form) 41.4 GPa
Shear modulus
(α form) 16.3 GPa
Bulk modulus
(α form) 31.8 GPa
Poisson ratio
−1
·K
−1
(α form) 0.281
Vickers hardness
343 MPa
Brinell hardness
265 MPa
CAS registry number
7440-00-8 Most-stable isotopes
Main article: Isotopes of neodymium iso
NA
half-life
142
27.2%
142
143
12.2%
143
144
23.8%
2.29×1015y
145
8.3%
145
146
17.2%
146
148
5.7%
148
150
5.6%
6.7×1018y
Nd Nd Nd Nd Nd Nd Nd
DM
DE (MeV)
DP
Nd is stable with 82 neutron Nd is stable with 83 neutron α
1.905
140
3.367
150
Ce
Nd is stable with 85 neutron Nd is stable with 86 neutron Nd is stable with 88 neutron β−β− References
Sm
[1]
Neodymium Neodymium (pronounced /ˌniː.ɵˈdɪmiəm/) is a chemical element with the symbol Nd and atomic number 60. It is a soft silvery metal which tarnishes in air. Neodymium was discovered in 1885. It is not found in nature in pure form; its various compounds are present in trace amounts in minerals monazite and bastnäsite. Neodymium has several important applications: it is a constituent of neodymium magnets, which are widely used in motors, loudspeakers and numerous appliances. Neodymium is a popular additive in glass, giving it a characteristic reddish-purple color; this glass is used in lasers emitting infrared light with the wavelength of 1.064 micrometers.
Characteristics Physical Neodymium, a rare earth metal, is present in mischmetal to the extent of about 18%. The metal has a bright, silvery metallic luster; however, as one of the more reactive rare earth (lanthanide) metals, it quickly oxidizes in air. The oxide layer then falls off, which exposes the metal to further oxidation. Neodymium exists in two allotropic forms, with a transformation from a double hexagonal to a body-centered cubic structure taking place at 863 °C.[2]
Chemical Neodymium metal tarnishes slowly in air and burns readily at 150 °C to form neodymium(III) oxide: 4 Nd + 3 O2 → 2 Nd2O3 Neodymium is quite electropositive and reacts slowly with cold water and quite quickly with hot water to form neodymium hydroxide: 2 Nd (s) + 6 H2O (g) → 2 Nd(OH)3 (aq) + 3 H2 (g) Neodymium metal reacts with all the halogens: 2 Nd (s) + 3 F2 (g) → 2 NdF3 (s) [violet] 2 Nd (s) + 3 Cl2 (g) → 2 NdCl3 (s) [mauve] 2 Nd (s) + 3 Br2 (g) → 2 NdBr3 (s) [violet] 2 Nd (s) + 3 I2 (g) → 2 NdI3 (s) [green] Neodymium dissolves readily in dilute sulphuric acid to form solutions containing the lilac Nd(III) ions, which exist as a [Nd(OH2)9]3+ complexes:[3] 2 Nd (s) + 3 H2SO4 (aq) → 2 Nd3+ (aq) + 3 SO2−4 (aq) + 3 H2 (g)
3
Neodymium
Compounds Neodymium compounds include • • • •
Halides: NdF3, NdCl3, NdBr3, NdI3 Oxides: Nd2O3 Sulfides: NdS, Nd2S3 Nitrides: NdN
Isotopes Naturally occurring neodymium is composed of 5 stable isotopes, 142Nd, 143Nd, 145Nd, 146 Nd and 148Nd, with 142Nd being the most abundant (27.2% natural abundance), and 2 radioisotopes, 144Nd and 150Nd. In all, 31 radioisotopes of Neodymium have been characterized up to now, with the most stable being naturally occurring isotopes 144Nd (alpha decay, a half-life (T½) of 2.29×1015 years) and 150Nd (double beta decay, T½ of 7×1018 years). All of the remaining radioactive isotopes have half-lives that are less than 11 days, and the majority of these have half-lives that are less than 70 seconds. This element also has 13 known meta states with the most stable being 139mNd (T½ 5.5 hours), 135mNd (T½ 5.5 minutes) and 133m1Nd (T½ ~70 seconds). The primary decay modes before the most abundant stable isotope, 142Nd, are electron capture and positron decay, and the primary mode after is beta minus decay. The primary decay products before 142Nd are element Pr (praseodymium) isotopes and the primary products after are element Pm (promethium) isotopes.
History Neodymium was discovered by Baron Carl Auer von Welsbach, an Austrian chemist, in Vienna in 1885. He separated neodymium, as well as the element praseodymium, from a material known as didymium by means of fractional crystallization of the double ammonium nitrate tetrahydrates from nitric acid, while following the separation by spectroscopic analysis; however, it was not isolated in relatively pure form until 1925. The name neodymium is derived from the Greek words neos, new, and didymos, twin. Neodymium is frequently misspelled as neodynium.[4] Double nitrate crystallization was the means of commercial neodymium purification until the 1950s. Lindsay Chemical Division was the first to commercialize large-scale ion-exchange purification of neodymium. Starting in the 1950s, high purity (e.g. 99+%) neodymium was primarily obtained through an ion exchange process from monazite, a mineral rich in rare earth elements. The metal itself is obtained through electrolysis of its halide salts. Currently, most neodymium is extracted from bastnaesite, (Ce,La,Nd,Pr)CO3F, and purified by solvent extraction. Ion-exchange purification is reserved for preparing the highest purities (typically >99.99 %). The evolving technology, and improved purity of commercially available neodymium oxide, was reflected in the appearance of neodymium glass made therefrom that resides in collections today. Early neodymium glass made in the 1930s, have a more reddish or orange tinge than modern versions, which are more cleanly purple, due to the difficulties in removing the last traces of praseodymium when the fractional crystallization technology had to be relied on.
4
Neodymium
Occurrence and production Neodymium is never found in nature as the free element; rather, it occurs in ores such as monazite and bastnäsite that contain small amounts of all the rare earth metals. The main mining areas are China, United States, Brazil, India, Sri Lanka and Australia; and reserves of neodymium are estimated as about 8 million tonnes. Although it belongs to "rare earth metals," neodymium is not rare at all - its abundance in the Bastnäsite Earth crust is about 38 mg/kg, which is the second among rare-earth elements after cerium. The world production of neodymium is about 7,000 tonnes per year.[4]
Applications • Neodymium magnets are the strongest permanent magnets known - Nd2Fe14B. These magnets are cheaper, lighter, and stronger than samarium-cobalt magnets. Neodymium magnets appear in products such as microphones, professional loudspeakers, in-ear headphones, guitar and bass guitar pick-ups Neodymium magnet on a bracket from and computer hard disks where low mass, small a hard drive. volume, or strong magnetic fields are required. Neodymium magnet electric motors have also been responsible for the development of purely electrical model aircraft within the first decade of the 21st century, to the point that these are displacing internal combustion powered models internationally. • Neodymium is a component of didymium used for coloring glass to make welder's and glass-blower's goggles. The sharp absorption bands obliterate the strong sodium emission at 589 nm. • Neodymium has an unusually large specific heat capacity at liquid-helium temperatures, so is useful in cryocoolers • Neodymium lamps are incandescent lamps containing neodymium in the glass to filter out yellow light, resulting in a whiter light more like sunlight • Neodymium colors glass in delicate shades ranging from pure violet through wine-red and warm grey. Light transmitted through such glass shows unusually sharp absorption bands; the glass is used in astronomical work to produce sharp bands by which spectral lines may be calibrated. Neodymium is also used to remove the green colour caused by iron contaminants from glass. • Neodymium salts are used as a colourant for enamels. • Probably because of similarities to Ca2+, Nd3+ has been reported[5] to promote plant growth. Rare earth element compounds are frequently used in China as fertilizer. • Samarium-neodymium dating is useful for determining the age relationships of rocks[6] and meteorites. • Size and strength of volcanic eruption can be predicted by scanning for neodymium isotopes. Small and large volcanic eruptions produce lava with different neodymium
5
Neodymium
6
isotope composition. From the composition of isotopes, scientists predict how big the coming eruption will be, and use this information to warn residents of the intensity of the eruption. • Certain transparent materials with a small concentration of neodymium ions can be used in lasers as gain media for infrared wavelengths (1054-1064 nm), e.g. Nd:YAG (yttrium aluminium garnet), Nd:YLF (yttrium lithium fluoride), Nd:YVO4 (yttrium orthovanadate), and Nd:glass. The current laser at the UK Atomic Weapons Establishment (AWE), the HELEN (High Energy Laser Embodying Neodymium) 1-terawatt neodymium-glass laser, can access the midpoints of pressure and temperature regions and is used to acquire data for modeling on how density, temperature and pressure interact inside warheads. HELEN can create plasmas of around 106 K, from which opacity and transmission of radiation are measured.[7]
Neodymium glass Neodymium
glass
(Nd:glass)
is
produced by the inclusion of neodymium oxide (Nd2O3) in the glass melt. In daylight or incandescent light neodymium glass appears lavender, but it appears pale blue under fluorescent lighting. Neodymium glass solid-state lasers
Neodymium doped glass slabs used in extremely powerful lasers for inertial confinement fusion.
are used in extremely high power (terawatt scale), high energy (megajoules) multiple beam systems for inertial confinement fusion. Nd:glass lasers are usually frequency tripled to the third harmonic at 351 nm in laser fusion
devices. Neodymium glass is becoming widely used in incandescent light bulbs, to provide a more "natural" light. It has been patented for use in automobile rear-view mirrors, to reduce the glare at night. The first commercial use of purified neodymium was in glass coloration, starting with experiments by Leo Moser in November 1927. The resulting "Alexandrite" glass remains a signature color of the Moser glassworks to this day. Neodymium glass was widely emulated in the early 1930s by American glasshouses, most notably Heisey, Fostoria ("wisteria"), Cambridge ("heatherbloom"), and Steuben ("wisteria"), and elsewhere (e.g. Lalique, in France, or Murano). Tiffin's "twilight" remained in production from about 1950 to about 1980.[8] Current sources include glassmakers in the Czech Republic, the United States, and China. The sharp absorption bands of neodymium cause the glass color to change under different lighting conditions, being reddish-purple under daylight or yellow incandescent light, but blue under white fluorescent lighting, or greenish under trichromatic lighting. This
Neodymium color-change phenomenon is highly prized by collectors. In combination with gold or selenium, beautiful red colors result. Since neodymium coloration depends upon "forbidden" f-f transitions deep within the atom, there is relatively little influence on the color from the chemical environment, so the color is impervious to the thermal history of the glass. However, for the best color, iron-containing impurities need to be minimized in the silica used to make the glass. The same "forbiddenness" of the f-f transitions makes rare-earth colorants less intense than those provided by most d-transition elements, so more has to be used in a glass to achieve the desired color intensity. The original Moser recipe used about 5% of neodymium oxide in the glass melt, a sufficient quantity such that Moser referred to these as being "rare earth doped" glasses. Being a strong base, that level of neodymium would have affected the melting properties of the glass, and the lime content of the glass might have had to be adjusted accordingly.[9]
Precautions Neodymium metal dust is a combustion and explosion hazard. Neodymium compounds, like all rare earth metals, are of low to moderate toxicity; however its toxicity has not been thoroughly investigated. Neodymium dust and salts are very irritating to the eyes and mucous membranes, and moderately irritating to skin. Breathing the dust can cause lung embolisms, and accumulated exposure damages the liver. Neodymium also acts as an anticoagulant, especially when given intravenously.[4] Neodymium magnets have been tested for medical uses such as magnetic braces and bone repair, but biocompatibility issues have prevented widespread application. Commercially available magnets made from Neodymium are exceptionally strong, and can attract each other from large distances. If not handled carefully, they could come together very quickly and forcefully, causing injuries. For example, a person lost part of his finger when two magnets he was using snapped together from 50 cm away.[10] Another danger is when two such magnets snap together, the force of the hit often causes them to shatter, sending sharp pieces flying around, potentially causing serious injuries.[4]
See also • Neodymium magnet (NIB or Nd2Fe14B)
Books • "The Industrial Chemistry of the Lanthanons, Yttrium, Thorium and Uranium", by R.J. Callow, Pergamon Press 1967. • Lindsay Chemical Division, American Potash and Chemical Corporation, Price List, 1960. • "Chemistry of the Lanthanons", by R.C. Vickery, Butterworths 1953.
7
Neodymium
8
External links • USGS Rare Earth Commodity Summary 2006 • WebElements.com – Neodymium [12] • It's Elemental – Neodymium [13]
[11]
pnb:میمئاڈوین
References [1] Gschneidner, K.A., and Eyring, L., Handbook on the Physics and Chemistry of Rare Earths, North Holland Publishing Co., Amsterdam, 1978. [2] C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304814. [3] " Chemical reactions of Neodymium (https:/ / www. webelements. com/ neodymium/ chemistry. html)". Webelements. . Retrieved 2009-06-06. [4] John Emsley (2003). Nature's building blocks: an A-Z guide to the elements (http:/ / books. google. co. jp/ books?id=j-Xu07p3cKwC). Oxford University Press. pp. 268–270. ISBN 0198503407. . [5] Y. Wei et al. "The Effect of Neodymium (Nd3+) on Some Physiological Activities in Oilseed Rape during Calcium (Ca2+) Starvation" 10th International Rapeceed Congress (http:/ / www. regional. org. au/ au/ gcirc/ 2/ 399. htm) [6] " Team finds Earth's 'oldest rocks' (http:/ / news. bbc. co. uk/ 2/ hi/ science/ nature/ 7639024. stm)". BBC news. . Retrieved 2009-06-06. [7] Norman, Michael J.; Andrew, James E.; Bett, Thomas H.; Clifford, Roger K.; England, John E.; Hopps, Nicholas W.; Parker, Kenneth W.; Porter, Kenneth; Stevenson, Mark (2002). "Multipass Reconfiguration of the HELEN Nd:Glass Laser at the Atomic Weapons Establishment". Applied Optics 41: 3497. doi: 10.1364/AO.41.003497 (http:/ / dx. doi. org/ 10. 1364/ AO. 41. 003497). [8] " Chameleon Glass Changes Color (http:/ / coloradosprings. yourhub. com/ CrippleCreekTellerCounty/ Stories/ Arts/ Story~443258. aspx)". . Retrieved 2009-06-06. [9] Charles Bray (2001). Dictionary of glass: materials and techniques (http:/ / books. google. com/ books?id=KbZkxDyeG18C& pg=PA102). University of Pennsylvania Press. p. 102. ISBN 081223619X. . [10] Swain, Frank (March 6, 2009). " How to remove a finger with two super magnets (http:/ / scienceblogs. com/ sciencepunk/ 2009/ 03/ how_to_remove_a_finger_with_tw. php)". Seed Media Group LLC. . Retrieved 2009-06-28. [11] http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ rare_earths/ [12] http:/ / www. webelements. com/ webelements/ elements/ text/ Nd/ index. html [13] http:/ / education. jlab. org/ itselemental/ ele060. html
Article Sources and Contributors
Article Sources and Contributors Neodymium Source: http://en.wikipedia.org/w/index.php?oldid=308825205 Contributors: A Mom, Ahoerstemeier, AlimanRuna, Ann Stouter, Archimerged, Arkuat, Awsoma, Beetstra, Benbest, Black Kite, BlueEarth, Bobo192, Borislav Dopudja, Breno, Bryan Derksen, Cacycle, Carnildo, Chetvorno, Chowbok, Chrisbolt, Cojoco, Comfychaos, Conversion script, CoolMike, Darrien, Davehi1, David Latapie, Deglr6328, Dennis Brown, DrBob, Dwmyers, Edgar181, El C, Emperorbma, Eranb, Eras-mus, Evil saltine, Excirial, Femto, Fivemack, Foobar, G.v.kolbe, Giftlite, Gr8white, Grendelkhan, Hadal, Hak-kâ-ngìn, Hankwang, Hannibal, Hellno2, Hibernian, HoodedMan, Icairns, Irdepesca572, Iridescent, Jaan513, Jacobst, Jaraalbe, JesseW, Joanjoc, John, Jose77, Kimse, Kwamikagami, LA2, LAMARCTRASK, LMB, LarryMorseDCOhio, Lenilucho, Looxix, Maniwar, Marc Venot, Materialscientist, Mav, Mboverload, Mentifisto, Minesweeper, Mmm, MorrisRob, Mysid, Nergaal, Plexust, Poolkris, Pras, RG2, RPaschotta, RTC, RandomCritic, Remember, Reza kalani, Roberta F., Rossheth, Roy W. Wright, RoyBoy, Rursus, Sam Hocevar, Saperaud, Schneelocke, Sengkang, Sfuerst, Sl, Stephan Leeds, Stephenb, Stevey7788, Stifynsemons, Stone, Suisui, Tagishsimon, Tetracube, Thrc, Thricecube, V1adis1av, Vsmith, Walkerma, Warut, WereSpielChequers, Yekrats, Yyy, Zanudaaa, Zoicon5, 155 anonymous edits
Image Sources, Licenses and Contributors image:Nd-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Nd-TableImage.png License: GNU Free Documentation License Contributors: user:schneelocke Image: Neodym 1-crop.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Neodym_1-crop.jpg License: unknown Contributors: User:Materialscientist File:Bastnaesite - Kischtimsk, Ural.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Bastnaesite_-_Kischtimsk,_Ural.jpg License: unknown Contributors: User:Ra'ike Image:Neodymag.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Neodymag.jpg License: Public Domain Contributors: Bloodshedder Image:Laser glass slabs.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Laser_glass_slabs.jpg License: Public Domain Contributors: Science and Technology Review, LLNLameee
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9
Promethium
1
Promethium 61
neodymium ← promethium → samarium
↑
Pm ↓
Np Periodic Table - Extended Periodic Table
General Name, Symbol, Number
promethium, Pm, 61
Element category
lanthanoids
Group, Period, Block
n/a, 6, f
Appearance
metallic
Standard atomic weight
[145](0) g·mol
Electron configuration
[Xe] 4f 6s
Electrons per shell
−1
5
2
2, 8, 18, 23, 8, 2 Physical properties
Phase
solid
Density (near r.t.)
7.26 g·cm
Melting point
1315 K (1042 °C, 1908 °F)
Boiling point
3273 K (3000 °C, 5432 °F)
Heat of fusion
7.13 kJ·mol−1
Heat of vaporization
289 kJ·mol−1
−3
Atomic properties Crystal structure Oxidation states Electronegativity Ionization energies (more)
hexagonal 3 (mildly basic oxide) ? 1.13 (Pauling scale) 1st: 540 kJ·mol−1 2nd: 1050 kJ·mol−1 3rd: 2150 kJ·mol−1
Atomic radius
183 pm
Covalent radius
199 pm Miscellaneous
Promethium
2 [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(r.t.) est. 0.75 µΩ·m
Thermal conductivity
(300 K) 17.9 W·m−1·K−1
Thermal expansion
(r.t.) (α, poly) est. 11 µm/(m·K)
Young's modulus
(α form) est. 46 GPa
Shear modulus
(α form) est. 18 GPa
Bulk modulus
(α form) est. 33 GPa
Poisson ratio
(α form) est. 0.28
CAS registry number
7440-12-2 Most-stable isotopes
Main article: Isotopes of promethium iso
NA trace
145
Pm
146
Pm
syn
half-life 17.7 y 5.53 y
DM ε
Pm
syn
2.6234 y
DP
0.163
145
1.472
146
−
1.542
146
−
0.224
147
ε β
147
DE (MeV)
β
Nd Nd Sm Sm
References
Promethium (pronounced /prɵˈmiːθiəm/) is a chemical element with the symbol Pm and atomic number 61. It is notable for being the only other exclusively radioactive element besides technetium which is followed by chemical elements that have stable isotopes.
Characteristics Physical Promethium's longest lived isotope 145Pm is a soft beta emitter with a half-life of 17.7 years. It does not emit gamma rays, but beta particles impinging on elements of high atomic numbers can generate X-rays, and a sample of 145Pm does produce some such soft X-ray radiation in addition to beta particles. Pure promethium exists in two allotropic forms, and its chemistry is similar to other lanthanides. Promethium salts luminesce in the dark with a pale blue or greenish glow, due to their high radioactivity. Promethium can be found in traces in some uranium ores, as a fission product. Newly made promethium is also seen in the spectra of some stars.
Promethium
Chemical Promethium metal tarnishes slowly in air and burns readily at 150 °C to form promethium(III) oxide: 4 Pm + 3 O2 → 2 Pm2O3 Promethium is quite electropositive and reacts slowly with cold water and quite quickly with hot water to form promethium hydroxide: 2 Pm (s) + 6 H2O (g) → 2 Pm(OH)3 (aq) + 3 H2 (g) Promethium metal reacts with all the halogens: 2 Pm (s) + 3 F2 (g) → 2 PmF3 (s) 2 Pm (s) + 3 Cl2 (g) → 2 PmCl3 (s) 2 Pm (s) + 3 Br2 (g) → 2 PmBr3 (s) 2 Pm (s) + 3 I2 (g) → 2 PmI3 (s) Promethium dissolves readily in dilute sulphuric acid to form solutions containing the pink Pm(III) ions, which exist as a [Pm(OH2)9]3+ complexes:[2] 2 Pm(s) + 3 H2SO4 (aq) → 2 Pm3+ (aq) + 3 SO2−4 (aq) + 3 H2 (g)
Compounds Promethium compounds include: • Chlorides • PmCl3 • Bromides • PmBr3 • Oxides • Pm2O3
Isotopes Thirty-six radioisotopes of promethium have been characterized, with the most stable being 145 Pm with a half-life of 17.7 years, 146Pm with a half-life of 5.53 years, and 147Pm with a half-life of 2.6234 years. All of the remaining radioactive isotopes have half-lives that are less than 364 days, and the majority of these have half lives that are less than 27 seconds. This element also has 11 meta states with the most stable being 148Pmm (T½ 41.29 days), 152 Pmm2 (T½ 13.8 minutes) and 152Pmm (T½ 7.52 minutes). The isotopes of promethium range in atomic weight from 127.9482600 u (128Pm) to 162.9535200 u (163Pm). The primary decay mode before the longest-lived isotope, 145Pm, is electron capture, and the primary mode after is beta minus decay. The primary decay products before 145Pm are neodymium (Nd) isotopes and the primary products after are samarium (Sm) isotopes. Along with technetium, promethium is one of only two elements with atomic number less than 83 that have only unstable isotopes, which is a rarely occurring effect of the liquid drop model and stabilities of neighbor element isotopes.
3
Promethium
4
History The existence of promethium was first predicted by Bohuslav Brauner in 1902; this prediction was supported in 1914 by Henry Moseley who, having discovered that atomic number was an experimentally measurable property of elements, found that no known element had atomic number 61. Several groups claimed to have produced the element, but they could not confirm their discoveries because of the difficulty of separating promethium from other elements. Promethium was first produced and proved to exist at Oak Ridge National Laboratory (ORNL) in 1945 by Jacob A. Marinsky, Lawrence E. Glendenin and Charles D. Coryell by separation and analysis of the fission products of uranium fuel irradiated in the Graphite Reactor; however, being too busy with defense-related research during World War II, they did not announce their discovery until 1947.[3] The name promethium is derived from Prometheus, the Titan, in Greek mythology, who stole the fire from Mount Olympus and brought it down to mankind. The name was suggested by Grace Mary Coryell, Charles Coryell's wife, who felt that they were stealing fire from the gods. In 1963, ion-exchange methods were used at ORNL to prepare about ten grams of promethium from nuclear reactor fuel processing wastes. Today, promethium is still recovered from the byproducts of uranium fission; it can also be produced by bombarding 146Nd with neutrons, turning it into 147Nd which decays into 147 Pm through beta decay with a half-life of 11 days.
Occurrence Promethium can be formed in nature as a product of spontaneous fission of uranium-238 and alpha decay of europium-151. Only trace amounts can be found in naturally occurring ores: a sample of pitchblende has been found to contain promethium at a concentration of four parts per quintillion (1018) by mass.[4] It was calculated that the equilibrium mass of promethium in the earth's crust is about 560 g due to uranium fission and about 12 g due to the recently observed alpha decay of europium-151.[5]
Pitchblende
Promethium has also been identified in the spectrum of the star HR 465 in Andromeda, and possibly HD 101065 (Przybylski's star) and HD 965.[6]
Applications Uses for promethium include: • As a beta radiation source for thickness gauges. • As a light source for signals that require reliable, independent operation (using phosphor to absorb the beta radiation and produce light). • In a nuclear battery in which cells convert the beta emissions into electric current, yielding a useful life of about five years, using Pm-147. • Promethium(III) chloride (PmCl3) mixed with zinc sulfide (ZnS) was used for a time as a major luminous paint for watches after radium was discontinued. This mixture is still occasionally used for some luminous paint applications (though most such uses with
Promethium
5
radioactive materials have switched to tritium for safety reasons). • Promethium has possible future uses in portable X-ray sources, and as auxiliary heat or power sources for space probes and satellites (although the alpha emitter plutonium-238 has become standard for most space-exploration related uses – see Radioisotope thermoelectric generators).
Precautions Promethium must be handled with great care because of its high radioactivity. In particular, promethium can emit X-rays during its beta decay. Its half-life is less than that of plutonium-239 by a factor of about 1350, and its biological toxicity is correspondingly higher. Promethium has no biological role.
External links • WebElements.com – Promethium • It's Elemental – Promethium [8]
[7]
pnb:میھتمورپ
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] " Chemical reactions of Promethium (https:/ / www. webelements. com/ promethium/ chemistry. html)". Webelements. . Retrieved 2009-06-06. [3] " Discovery of Promethium (http:/ / www. ornl. gov/ info/ ornlreview/ v36_1_03/ article_02. shtml)". ORNL Review 36 (1). 2003. . Retrieved 2006-09-17. [4] Attrep, Moses, Jr.; and P. K. Kuroda (May 1968). "Promethium in pitchblende". Journal of Inorganic and Nuclear Chemistry 30 (3): 699–703. doi: 10.1016/0022-1902(68)80427-0 (http:/ / dx. doi. org/ 10. 1016/ 0022-1902(68)80427-0). [5] P. Belli, R. Bernabei, F. Cappella, R. Cerulli, C.J. Dai, F.A. Danevich, A. d’Angelo, A. Incicchitti, V.V. Kobychev, S.S. Nagorny, S. Nisi, F. Nozzoli, D. Prosperi, V.I. Tretyak, S.S. Yurchenko (2007). "Search for α decay of natural Europium". Nuclear Physics A 789: 15–29. doi: 10.1016/j.nuclphysa.2007.03.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2007. 03. 001). [6] C. R. Cowley, W. P. Bidelman, S. Hubrig, G. Mathys, and D. J. Bord (2004). "On the possible presence of promethium in the spectra of HD 101065 (Przybylski's star) and HD 965". Astronomy & Astrophysics 419: 1087–1093. doi: 10.1051/0004-6361:20035726 (http:/ / dx. doi. org/ 10. 1051/ 0004-6361:20035726). [7] http:/ / www. webelements. com/ webelements/ elements/ text/ Pm/ index. html [8] http:/ / education. jlab. org/ itselemental/ ele061. html
Article Sources and Contributors
Article Sources and Contributors Promethium Source: http://en.wikipedia.org/w/index.php?oldid=308865583 Contributors: 2over0, AP1787, Aadal, Ahoerstemeier, AlimanRuna, Andre Engels, Anthony Appleyard, Archimerged, Arkuat, AstroHurricane001, Benbest, Benjiboi, Bettia, Bkell, BlueEarth, Borislav Dopudja, Bryan Derksen, Carnildo, Conversion script, Croftfoothen, Darklilac, Darrien, David Latapie, Deor, Donarreiskoffer, DragonflySixtyseven, Dyslexic Q-Thief, Edgar181, El C, Emperorbma, Encyclopedia77, Evil saltine, Femto, Fivemack, Gazzarrr, Gordonmichaels, Hellbus, Herbee, Hqb, IForgotToEatBreakFast, II MusLiM HyBRiD II, Icairns, Ideyal, IronGargoyle, JNW, JWB, JYolkowski, Jaan513, JackofOz, Jaraalbe, Jeneralist, Joanjoc, Jose77, Karl-Henner, Karlhahn, Keenan Pepper, Kostisl, Kwamikagami, LA2, Looxix, Lord Pistachio, MC10, Marc Venot, Materialscientist, Mav, Maxamegalon2000, Michbich, Minesweeper, Mmm, Mortdefides, NHRHS2010, Nergaal, Nick Y., NoamTene, Orlady, Pak21, Pakaran, Paul1953h, PierreAbbat, Poolkris, Pras, Pyfan, Red Director, Remember, Reyk, Roberta F., Romanm, Romeu, Rrburke, Rursus, Saperaud, Sbharris, Schneelocke, Segin, Sengkang, Sl, Stack, Stifynsemons, Stone, Stratocracy, Sunnyoraish, Svante, Svlad Jelly, Tagishsimon, TerraFrost, Tetracube, Thinghy, V1adis1av, Van helsing, Vsmith, Vuerqex, Walkerma, Warut, Yekrats, Yyy, 102 anonymous edits
Image Sources, Licenses and Contributors image:Pm-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Pm-TableImage.png License: GNU Free Documentation License Contributors: user:schneelocke File:Pitchblende schlema-alberoda.JPG Source: http://en.wikipedia.org/w/index.php?title=File:Pitchblende_schlema-alberoda.JPG License: unknown Contributors: User:Geomartin
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6
Samarium
1
Samarium 62
promethium ← samarium → europium
↑
Sm ↓
Pu Periodic Table - Extended Periodic Table
General Name, Symbol, Number
samarium, Sm, 62
Element category
lanthanides
Group, Period, Block
n/a, 6, f
Appearance
silvery white
Standard atomic weight
150.36(2) g·mol
Electron configuration
[Xe] 6s 4f
−1
2
Electrons per shell
6
2, 8, 18, 24, 8, 2 Physical properties
Phase
solid
Density (near r.t.)
7.52 g·cm
Liquid density at m.p.
7.16 g·cm−3
Melting point
1345 K (1072 °C, 1962 °F)
Boiling point
2067 K (1794 °C, 3261 °F)
Heat of fusion
8.62 kJ·mol−1
Heat of vaporization
165 kJ·mol−1
Specific heat capacity
(25 °C) 29.54 J·mol−1·K−1
−3
Vapor pressure P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1001
1106
1240
(1421)
(1675)
(2061)
Atomic properties Crystal structure
rhombohedral
Samarium
2
Oxidation states
3, 2 (mildly basic oxide)
Electronegativity
1.17 (Pauling scale)
Ionization energies (more)
1st: 544.5 kJ·mol−1 2nd: 1070 kJ·mol−1 3rd: 2260 kJ·mol−1
Atomic radius
180 pm
Covalent radius
198±8 pm Miscellaneous [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(r.t.) (α, poly) 0.940 µΩ·m
Thermal conductivity
(300 K) 13.3 W·m
Thermal expansion
(r.t.) (α, poly) 12.7 µm/(m·K)
Speed of sound (thin rod)
(20 °C) 2130 m/s
Young's modulus
(α form) 49.7 GPa
Shear modulus
(α form) 19.5 GPa
Bulk modulus
(α form) 37.8 GPa
Poisson ratio
−1
·K
−1
(α form) 0.274
Vickers hardness
412 MPa
Brinell hardness
441 MPa
CAS registry number
7440-19-9 Most-stable isotopes
Main article: Isotopes of samarium iso 144
Sm
146
Sm
NA 3.07% syn
half-life
DM
DE (MeV)
DP
144
Sm is stable with 82 neutron
1.03×108y
α
2.529
142
Nd
147
14.99%
1.06×1011y
α
2.310
143
148
11.24%
7×1015y
α
1.986
144
149
13.82%
α
1.870
145
150
7.38%
150
152
26.75%
152
154
22.75%
154
Sm Sm Sm Sm Sm Sm
>2×1015 y
Sm is stable with 88 neutron Sm is stable with 90 neutron Sm is stable with 92 neutron References
Nd Nd Nd
Samarium Samarium (pronounced /səˈmɛəriəm/) is a chemical element with the symbol Sm and atomic number 62.
Characteristics Physical Samarium is a rare earth metal, with a bright silver luster. Three crystal modifications of the metal also exist, with transformations at 734 and 922 °C, making it polymorphic. Individual samarium atoms can be isolated by encaspulating them into fullerene molecules.[2]
Chemical Samarium oxidizes in air and ignites at 150 °C. Even with long-term storage under mineral oil, samarium is gradually oxidized, with a grayish-yellow powder of the oxide-hydroxide being formed. The metallic appearance of a sample can be preserved by sealing it under an inert gas such as argon. Samarium is quite electropositive and reacts slowly with cold water and quite quickly with hot water to form samarium hydroxide: 2 Sm (s) + 6 H2O (g) → 2 Sm(OH)3 (aq) + 3 H2 (g) Samarium metal reacts with all the halogens: 2 Sm (s) + 3 F2 (g) → 2 SmF3 (s) [white] 2 Sm (s) + 3 Cl2 (g) → 2 SmCl3 (s) [yellow] 2 Sm (s) + 3 Br2 (g) → 2 SmBr3 (s) [yellow] 2 Sm (s) + 3 I2 (g) → 2 SmI3 (s) [orange] Samarium dissolves readily in dilute sulphuric acid to form solutions containing the pale green Sm(III) ions, which exist as a [Sm(OH2)9]3+ complexes:[3] 2 Sm (s) + 3 H2SO4 (aq) → 2 Sm3+ (aq) + 3 SO2−4 (aq) + 3 H2 (g)
Compounds Compounds of Samarium include: • • • • • • • •
Fluorides: SmF2, SmF3 Chlorides: SmCl2, SmCl3 Bromides: SmBr2, SmBr3 Iodides: SmI2, SmI3 Oxides: Sm2O3 Sulfides: Sm2S3 Selenides: Sm2Se3 Tellurides: Sm2Te3
The most common oxidation state of samarium is +3, but +2 compounds are known too, such as SmI2.
3
Samarium
Isotopes Naturally occurring samarium is composed of four stable isotopes, 144Sm, 150Sm, 152Sm and 154Sm, and three extremely long-lived radioisotopes, 147Sm (1.06 × 1011y), 148Sm (7 × 1015y) and 149Sm (>2 × 1015y), with 152Sm being the most abundant (26.75% natural abundance). 151
Sm has a halflife of 90 years, and 145Sm has a halflife of 340 days. All of the remaining radioisotopes have half-lives that are less than 2 days, and the majority of these have half-lives that are less than 48 seconds. This element also has 5 meta states with the most stable being 141mSm (t½ 22.6 minutes), 143m1Sm (t½ 66 seconds) and 139mSm (t½ 10.7 seconds). The primary decay mode before the most abundant stable isotope, 152Sm, is electron capture, and the primary mode after is beta minus decay. The primary decay products before 152Sm are element Pm (promethium) isotopes, and the primary products after are element Eu (europium) isotopes. Natural Samarium has an activity of 128 Bq/g.
History Samarium was first discovered spectroscopically in 1853 by Swiss chemist Jean Charles Galissard de Marignac by its sharp absorption lines in didymium, and isolated in Paris in 1879 by French chemist Paul Émile Lecoq de Boisbaudran from the mineral samarskite ((Y,Ce,U,Fe)3(Nb,Ta,Ti)5O16). Although samarskite was first found in the Urals, by the late 1870s a new deposit had been located in North Carolina, and it was from that source that the samarium-bearing didymium had originated. The samarskite mineral was named after Vasili Samarsky-Bykhovets, the Chief of Staff (Colonel) of the Russian Corps of Mining Engineers in 1845–1861. The name of the element is derived from the name of the mineral, and thus traces back to the name Samarsky-Bykhovets. In this sense samarium was the first chemical element to be named after a living person. Prior to the advent of ion-exchange separation technology in the 1950s, samarium had no commercial uses in pure form. However, a by-product of the fractional crystallization purification of neodymium was a mixture of samarium and gadolinium that acquired the name of "Lindsay Mix" after the company that made it. This material is thought to have been used for nuclear control rods in some of the early nuclear reactors. Nowadays, a similar commodity product goes under the name of "Samarium-Europium-Gadolinium" concentrate (or SEG concentrate). This is prepared by solvent extraction from the mixed lanthanides extracted from bastnäsite (or monazite). Since the heavier lanthanides have the greater affinity for the solvent used, they are easily extracted from the bulk using relatively small proportions of solvent. Not all rare earth producers who process bastnäsite do so on large enough scale to continue onward with the separation of the components of SEG, which typically makes up only one or two percent of the original ore. Such producers will therefore be making SEG with a view to marketing it to the specialized processors. In this manner, the valuable europium content of the ore is rescued for use in phosphor manufacture. Samarium purification follows the removal of the europium. Currently, being in oversupply, samarium oxide is less expensive on a commercial scale than its relative abundance in the ore might suggest.
4
Samarium
Occurrence Samarium is never found free in nature, but, like other rare earth elements, is contained in many minerals, including monazite, bastnäsite and samarskite; monazite (in which it occurs up to an extent of 2.8%) and bastnäsite are also used as commercial sources. Misch metal containing about 1% of samarium has long been used, but it was not until recent years that relatively pure samarium has been isolated through ion exchange processes, solvent extraction techniques, and Samarskite electrochemical deposition. The metal is often prepared by electrolysis of a molten mixture of samarium(III) [1] chloride with sodium chloride or calcium chloride . Samarium can also be obtained by reducing its oxide with lanthanum.
Applications Uses of Samarium include: • Carbon-arc lighting for the motion picture industry (together with other rare earth metals). • CaF2 crystals for use in lasers. • As a neutron absorber in nuclear reactors. • For alloys. • For headphone magnets. • Samarium-cobalt magnets; SmCo5 and Sm2Co17 are used in making permanent magnet materials that have high resistance to demagnetization when compared to other permanent magnet materials. These materials have high coercivities and intrinsic coercivities. Samarium-cobalt combinations have recently found use in high-end magnetic pickups for guitars and related musical instruments. • Samarium(II) iodide is used as a reducing agent and coupling agent chemical reagent in organic synthesis, for example in the Barbier reaction.[4] • Samarium oxide is used in optical glass to absorb infrared light. • Samarium compounds act as sensitizers for phosphors excited in the infrared. • Samarium oxide is a catalyst for the dehydration and dehydrogenation of ethanol. • Samarium-neodymium dating is useful for determining the age relationships of rocks and meteorites. • Radioactive Samarium-153 is used in medicine to treat the severe pain associated with cancers that have spread to bone. The drug is called "Quadramet".[5]
5
Samarium
6
Precautions As with the other lanthanides, samarium compounds are of low to moderate toxicity, although their toxicity has not been investigated in detail.
Books • N. N. Greenwood, A. Earnshaw, Chemistry of the Elements, Pergamon Press, Oxford, UK, 1984.
External links • WebElements.com – Samarium [6] • It's Elemental – Samarium [7] • Reducing Agents > Samarium low valent
[8]
pnb:میراماس
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] Okazaki , T (2002). "Electronic and geometric structures of metallofullerene peapods". Physica B: Condensed Matter 323: 97. doi: 10.1016/S0921-4526(02)00991-2 (http:/ / dx. doi. org/ 10. 1016/ S0921-4526(02)00991-2). [3] " Chemical reactions of Samarium (https:/ / www. webelements. com/ samarium/ chemistry. html)". Webelements. . Retrieved 2009-06-06. [4] Cotton (2007). Advanced inorganic chemistry, 6th ed (http:/ / books. google. com/ books?id=U3MWRONWAmMC& pg=PA1128). Wiley-India. p. 1128. ISBN 8126513381. . [5] " Centerwatch About drug Quadramet (http:/ / www. centerwatch. com/ patient/ drugs/ dru267. html)". . Retrieved 2009-06-06. [6] http:/ / www. webelements. com/ webelements/ elements/ text/ Sm/ index. html [7] http:/ / education. jlab. org/ itselemental/ ele062. html [8] http:/ / www. organic-chemistry. org/ chemicals/ reductions/ samariumlowvalent. shtm
Article Sources and Contributors
Article Sources and Contributors Samarium Source: http://en.wikipedia.org/w/index.php?oldid=308594324 Contributors: ARHAPSTF, Ahoerstemeier, AlimanRuna, Archimerged, Arkuat, Arteitle, B.d.mills, Barneca, Benbest, Borislav Dopudja, Brian Huffman, Bryan Derksen, Bucketsofg, Cacycle, Calibas, Carnildo, ChrisNanji, Closedmouth, Conversion script, DMacks, Darrien, David Latapie, Debresser, Discospinster, Donarreiskoffer, Edgar181, Emperorbma, ErikvdL, Femto, Finn-Zoltan, Gdommett, Goudzovski, Grendelkhan, Hak-kâ-ngìn, Hankwang, Helge Skjeveland, Hqb, Icairns, Ideyal, IvanLanin, J.delanoy, JWB, Jaan513, Jaganath, Jaraalbe, Joanjoc, JohnCD, Jose77, Kalamkaar, Karl-Henner, Kostisl, Ktsquare, Kurykh, Kwamikagami, LA2, Lando5, LarryMorseDCOhio, Looxix, Madmedea, Mani1, Marc Venot, Mare Nostrum, Materialscientist, Mav, Maxis ftw, Mdf, Mike Rosoft, Minesweeper, Nergaal, Nlu, Notheruser, P-Squared1969, PeepP, Pharaoh of the Wizards, Plexust, Poolkris, Principalityofgalore, Prosfilaes, QaviArtilles, RTC, Remember, Roberta F., Romanm, RunOrDie, Sam Hocevar, Saperaud, Schewek, Schneelocke, Sengkang, Sfuerst, Sheitan, Shirulashem, Sl, Stifynsemons, Suisui, Svante, Tagishsimon, Tetracube, Tiddly Tom, Timeastor, Unreal128, V1adis1av, Vsmith, WFPM, Walkerma, Warut, Wiki alf, Xxanthippe, Yekrats, Yyy, 93 anonymous edits
Image Sources, Licenses and Contributors image:Sm-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Sm-TableImage.png License: GNU Free Documentation License Contributors: user:schneelocke Image: Samarium 1-cropflipped.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Samarium_1-cropflipped.jpg License: unknown Contributors: User:Materialscientist File:Samarskite.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Samarskite.jpg License: unknown Contributors: Photograph by Andrew Silver. Original uploader was Tillman at en.wikipedia
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7
Europium
1
Europium 63
samarium ← europium → gadolinium
↑
Eu ↓
Am Periodic Table - Extended Periodic Table
General Name, Symbol, Number
europium, Eu, 63
Element category
lanthanides
Group, Period, Block
n/a, 6, f
Appearance
silvery white
Standard atomic weight
151.964(1) g·mol
Electron configuration
[Xe] 4f 6s
−1
7
Electrons per shell
2
2, 8, 18, 25, 8, 2 Physical properties
Phase
solid
Density (near r.t.)
5.264 g·cm
Liquid density at m.p.
5.13 g·cm
Melting point
1099 K (826 °C, 1519 °F)
Boiling point
1802 K (1529 °C, 2784 °F)
Heat of fusion
9.21 kJ·mol−1
Heat of vaporization
176 kJ·mol−1
Specific heat capacity
(25 °C) 27.66 J·mol−1·K−1
−3
−3
Vapor pressure P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
863
957
1072
1234
1452
1796
Atomic properties Crystal structure
simple cubic (body centered)
Europium
2
Oxidation states
3, 2 (mildly basic oxide)
Electronegativity
? 1.2 (Pauling scale)
Ionization energies (more)
1st: 547.1 kJ·mol−1 2nd: 1085 kJ·mol−1 3rd: 2404 kJ·mol−1
Atomic radius
180 pm
Covalent radius
198±6 pm Miscellaneous [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(r.t.) (poly) 0.900 µΩ·m
Thermal conductivity
(300 K) est. 13.9 W·m
Thermal expansion
(r.t.) (poly) 35.0 µm/(m·K)
Young's modulus
18.2 GPa
Shear modulus
7.9 GPa
Bulk modulus
8.3 GPa
Poisson ratio
0.152
−1
Vickers hardness
−1
·K
167 MPa
CAS registry number
7440-53-1 Most-stable isotopes
Main article: Isotopes of europium iso 150
Eu
151
Eu
152
Eu
153
Eu
NA syn 47.8% syn
52.2%
half-life
DM
DE (MeV) 2.261
DP
36.9 y
ε
5×1018 y
α
13.516 y
ε
1.874
152
β−
1.819
152
150
Sm
147
Pm Sm Gd
153
Eu is stable with 90 neutron References
Europium (pronounced /jʊˈroʊpiəm/) is a chemical element with the symbol Eu and atomic number 63. It was named after the continent Europe.
Europium
3
Characteristics Physical It is about as hard as lead and quite ductile. Europium is a metal and becomes a superconductor under pressure 80 GPa at temperature 1.8 K.[2]
Chemical Europium is one of the most reactive of the rare earth elements; it rapidly oxidizes in air, and resembles calcium in its reaction with water; samples of the metal Dendritic sublimated Eu (~300 g; element in solid form, even when coated with a purity 99,998%) protective layer of mineral oil, are rarely shiny. Europium ignites in air at 150 °C to 180 °C to form europium(III) oxide: 4 Eu + 3 O2 → 2 Eu2O3 Europium is quite electropositive and reacts slowly with cold water and quite quickly with hot water to form europium hydroxide: 2 Eu (s) + 6 H2O (g) → 2 Eu(OH)3 (aq) + 3 H2 (g) Europium metal reacts with all the halogens: 2 Eu (s) + 3 F2 (g) → 2 EuF3 (s) [white] 2 Eu (s) + 3 Cl2 (g) → 2 EuCl3 (s) [yellow] 2 Eu (s) + 3 Br2 (g) → 2 EuBr3 (s) [gray] 2 Eu (s) + 3 I2 (g) → 2 EuI3 (s) Europium dissolves readily in dilute sulphuric acid to form solutions containing very pale pink Eu(III) ions, which exist as a [Eu(OH2)9]3+ complexes:[3] 2 Eu (s) + 3 H2SO4 (aq) → 2 Eu3+ (aq) + 3 SO (aq) + 3 H2 (g)
Compounds Europium compounds include: • • • • • • • • •
Fluorides: EuF2, EuF3 Chlorides: EuCl2, EuCl3 Bromides: EuBr2, EuBr3 Iodides: EuI2, EuI3 Oxides: EuO, Eu2O3, Eu3O4 Sulfides: EuS Selenides: EuSe Tellurides: EuTe Nitrides: EuN
Europium(II) compounds tend to predominate, in contrast to most lanthanides: (which generally form compounds with an oxidation state of +3). Europium(II) chemistry is very similar to barium(II) chemistry, as they have similar ionic radii. Divalent europium is a mild reducing agent, such that under atmospheric conditions, it is the trivalent form that predominates. Under anaerobic, and particularly under geothermal conditions, the divalent
Europium
4
form is sufficiently stable such that it tends to be incorporated into minerals of calcium and the other alkaline earths. This is the cause of the "negative europium anomaly", that depletes europium from being incorporated into the most usual light lanthanide minerals such as monazite, relative to the chondritic abundance. Bastnäsite tends to show less of a negative europium anomaly than monazite does, and hence is the major source of europium today. The accessible divalency of europium has always made it one of the easiest lanthanides to extract and purify, even when present, as it usually is, in low concentration.
Isotopes Naturally occurring europium is composed of 2 isotopes, 151Eu and 153Eu, with 153Eu being the most abundant (52.2% natural abundance). While 153Eu is stable, 151Eu was recently yr[4] (in reasonable
found to be unstable to alpha decay with half-life of
agreement with theoretical predictions), giving about 1 alpha decay per two minutes in every kilogram of natural europium. Besides natural radioisotope 151Eu, 35 artificial radioisotopes have been characterized, with the most stable being 150Eu with a half-life of 36.9 years, 152Eu with a half-life of 13.516 years, and 154Eu with a half-life of 8.593 years. All of the remaining radioactive isotopes have half-lives that are less than 4.7612 years, and the majority of these have half-lives that are less than 12.2 seconds. This element also has 8 meta states, with the most stable being 150mEu (T½=12.8 hours), 152m1Eu (T½=9.3116 hours) and 152m2Eu (T½=96 minutes). The primary decay mode before the most abundant stable isotope, 153Eu, is electron capture, and the primary mode after is beta minus decay. The primary decay products before 153Eu are isotopes of samarium (Sm) and the primary products after are isotopes of gadolinium (Gd).
Europium as a nuclear fission product Thermal neutron capture cross sections Isotope
151
152
153
154
155
Yield
~10
low
1580
>2.5
330
Barns
5900
12800
312
1340
3950
Eu
Eu
Eu
Eu
Eu
Medium-lived fission products Prop: Unit:
Yield %
t½ a
252
βγ
Kr 10.76
.2180
687
βγ
Cd 14.1
.0008
316 β
Sr 28.9
4.505
2826 β
Cs 30.23
6.337
1176
βγ
.00005
390
βγ
Eu
85
113m
βγ *
.0803
155
90
137
121m
Sn 43.9
4.76
Q* KeV
Europium
5 Sm 90
151
.5314
77 β
Europium is produced by nuclear fission, but the fission product yields of europium isotopes are low near the top of the mass range for fission products. Like other lanthanides, many isotopes, especially isotopes with odd mass numbers and neutron-poor isotopes like 152Eu, have high cross sections for neutron capture, often high enough to be neutron poisons. 151
Eu is the beta decay product of Sm-151, but since this has a long decay half-life and short mean time to neutron absorption, most 151Sm instead winds up as 152Sm. 152
Eu (half-life 13.516 years) and 154Eu (halflife 8.593 years) cannot be beta decay products because 152Sm and 154Sm are nonradioactive, but 154Eu is the only long-lived "shielded" nuclide, other than 134Cs, to have a fission yield of more than 2.5 parts per million fissions.[5] A larger amount of 154Eu will be produced by neutron activation of a significant portion of the nonradioactive153Eu; however, much of this will be further converted to 155 Eu. 155
Eu (halflife 4.7612 years) has a fission yield of 330 ppm for U-235 and thermal neutrons. Most will be transmuted to nonradioactive and nonabsorptive Gadolinium-156 by the end of fuel burnup. Overall, europium is overshadowed by Cs-137 and Sr-90 as a radiation hazard, and by samarium and others as a neutron poison.
History Europium was first found by Paul Émile Lecoq de Boisbaudran in 1890, who obtained basic fraction from samarium-gadolinium concentrates which had spectral lines not accounted for by samarium or gadolinium; however, the discovery of europium is generally credited to French chemist Eugène-Anatole Demarçay, who suspected samples of the recently discovered element samarium were contaminated with an unknown element in 1896 and who was able to isolate europium in 1901. When the europium-doped yttrium orthovanadate red phosphor was discovered in the early 1960s, and understood to be about to cause a revolution in the color television industry, there was a mad scramble for the limited supply of europium on hand among the monazite processors. (Typical europium content in monazite was about 0.05%.) Luckily, Molycorp, with its bastnäsite deposit at Mountain Pass, California, whose lanthanides had an unusually "rich" europium content of 0.1%, was about to come on-line and provide sufficient europium to sustain the industry. Prior to europium, the color-TV red phosphor was very weak, and the other phosphor colors had to be muted, to maintain color balance. With the brilliant red europium phosphor, it was no longer necessary to mute the other colors, and a much brighter color TV picture was the result. Europium has continued in use in the TV industry ever since, and, of course, also in computer monitors. Californian bastnäsite now faces stiff competition from Bayan Obo, China, with an even "richer" europium content of 0.2%. Frank Spedding, celebrated for his development of the ion-exchange technology that revolutionized the rare earth industry in the mid-1950s once related the story of how, in the 1930s, he was lecturing on the rare earths when an elderly gentleman approached him with an offer of a gift of several pounds of europium oxide. This was an unheard-of quantity at the time, and Spedding did not take the man seriously. However, a package duly arrived in the mail, containing several pounds of genuine europium oxide. The elderly gentleman had turned out to be Dr. McCoy who had
Europium
6
developed a famous method of europium purification involving redox chemistry.
Occurrence Europium is never found in nature as a free element; however, there are many minerals containing europium, with the most important sources being bastnäsite and monazite. Europium has also been identified in the spectra of the sun and certain stars. Depletion or enrichment of europium in minerals relative to other rare earth elements is known as the europium anomaly. Divalent europium in small amounts happens to be the activator of the bright blue fluorescence of some samples of the mineral fluorite (calcium difluoride). The Monazite most outstanding examples of this originated around Weardale, and adjacent parts of northern England, and indeed it was this fluorite that gave its name to the phenomenon of fluorescence, although it was not until much later that europium was discovered or determined to be the cause.
Production Europium is found in minerals xenotime, monazite, and bastnäsite. The first two are orthophosphate minerals LnPO4 (Ln denotes a mixture of all the lanthanoids except promethium which is vanishingly rare due to being radioactive) and the third is a fluoride carbonate LnCO3F. Lanthanoids with even atomic numbers are more common. The most common lanthanoids in these minerals are, in order, cerium, lanthanum, neodymium, and praseodymium. Monazite also contains thorium and yttrium, which makes handling difficult since thorium and its decomposition products are radioactive. For many purposes it is not particularly necessary to separate the metals, but if separation into individual metals is required, the process is complex. Initially, the metals are extracted as salts from the ores by extraction with sulfuric acid (H2SO4), hydrochloric acid (HCl), and sodium hydroxide (NaOH). Modern purification techniques for these lanthanoid salt mixtures are ingenious and involve selective complexation techniques, solvent extractions, and ion exchange chromatography.[6] Pure europium is available through the electrolysis of a mixture of molten EuCl3 and NaCl (or CaCl2) in a graphite cell which acts as cathode, using graphite as anode. The other product is chlorine gas.
Europium
7
Applications There are many commercial applications for europium metal: it has been used to dope some types of glass to make lasers, as well as for screening for Down syndrome and some other genetic diseases. Due to its amazing ability to absorb neutrons, it is also being studied for use in nuclear reactors. Europium oxide (Eu2O3) is widely used as a red phosphor in television sets and fluorescent lamps, and as an activator for yttrium-based phosphors. Whereas trivalent europium gives red phosphors, the luminescence of divalent europium depends on the host lattice, but tends to be on the blue side. The two europium phosphor classes Europium is one of the elements used (red and blue), combined with the yellow/green terbium to make the red color in CRT televisions. phosphors give "white" light, the color temperature of which can be varied by altering the proportion or specific composition of the individual phosphors. This is the phosphor system typically encountered in the helical fluorescent lightbulbs. Combining the same three classes is one way to make trichromatic systems in TV and computer screens. It is also being used as an agent for the manufacture of fluorescent glass. Europium fluorescence is used to interrogate biomolecular interactions in drug-discovery screens. It is also used in the anti-counterfeiting phosphors in Euro banknotes. [7] Europium is commonly included in trace element studies in geochemistry and petrology to understand the processes that form igneous rocks (rocks that cooled from magma or lava). The nature of the europium anomaly found is used to help reconstruct the relationships within a suite of igneous rocks.
Precautions The toxicity of europium compounds has not been fully investigated, but there are no clear indications that europium is highly toxic compared to other heavy metals. The metal dust presents a fire and explosion hazard. Europium has no known biological role.
See also • Europium anomaly
External links • WebElements.com – Europium • It's Elemental – Europium [9]
[8]
Europium
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] M. Debessai et al. (2009). "Pressure-Induced Superconducting State of Europium Metal at Low Temperatures". Phys. Rev. lett. 102: 197002. doi: 10.1103/PhysRevLett.102.197002 (http:/ / dx. doi. org/ 10. 1103/ PhysRevLett. 102. 197002). [3] " Chemical reactions of Europium (https:/ / www. webelements. com/ europium/ chemistry. html)". Webelements. . Retrieved 2009-06-06. [4] Search for α decay of natural Europium, P. Belli, R. Bernabei, F. Cappell, R. Cerulli, C.J. Dai, F.A. Danevich, A. d'Angelo, A. Incicchitti, V.V. Kobychev, S.S. Nagorny, S. Nisi, F. Nozzoli, D. Prosperi, V.I. Tretyak, and S.S. Yurchenko, Nucl. Phys. A 789, 15 (2007) doi: 10.1016/j.nuclphysa.2007.03.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2007. 03. 001) [5] ORNL Table of the Nuclides [6] Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds (http:/ / books. google. com/ books?id=Xqj-TTzkvTEC& pg=PA243). McGraw-Hill. pp. 294-295. ISBN 0070494398. . Retrieved 2009-06-06. [7] " Europium and the Euro (http:/ / www. smarterscience. com/ eurosandeuropium. html)". . Retrieved 2009-06-06. [8] http:/ / www. webelements. com/ webelements/ elements/ text/ Eu/ index. html [9] http:/ / education. jlab. org/ itselemental/ ele063. html
8
Article Sources and Contributors
Article Sources and Contributors Europium Source: http://en.wikipedia.org/w/index.php?oldid=308258278 Contributors: 2over0, AWeishaupt, Ahoerstemeier, AlimanRuna, Archimerged, Arkuat, B07, Benbest, Billytrousers, BlueEarth, Borislav Dopudja, Brian Huffman, Bryan Derksen, Carnildo, Chiu frederick, Chowbok, Christian List, Conversion script, Cryptic C62, Darrien, David Latapie, Deb, Deglr6328, Deli nk, Donarreiskoffer, Download, Eastlaw, Edgar181, Eleassar777, Element16, Emperorbma, Epbr123, Erik9, Femto, Frencheigh, Gail, Gene Nygaard, Grendelkhan, Gökhan, Hankwang, Helge Skjeveland, IRP, Icairns, Ideyal, Itub, IvanLanin, J.delanoy, JWB, Jaan513, Jake Wartenberg, Janche, Jaraalbe, Jkc0113, Joanjoc, Jojorocko, Jonathanischoice, Jose77, KFP, Kaihsu, Karelj, Karl-Henner, Kurykh, Kwamikagami, LA2, Latka, Looxix, Marc Venot, Materialscientist, Mav, McGeddon, Melaen, Minesweeper, Mortdefides, Neparis, Nergaal, Nsaa, PeepP, Phil Boswell, Picapica, Plexust, Poolkris, Psyche825, R N Talley, RTC, Remember, Rholton, Roberta F., Romanm, Ronhjones, Rune.welsh, SKREAM, Sakus, Sam Hocevar, Saperaud, Schneelocke, Sengkang, Signalhead, Sjö, Sl, Squids and Chips, Stephenb, Stifynsemons, Stone, Tagishsimon, Tetracube, Thingg, Tide rolls, Trevorash, Tukss, V1adis1av, VASANTH S.N., Van helsing, Vicki Rosenzweig, Vsmith, Walkerma, WereSpielChequers, Williamborg, Xy7, Yekrats, Yyy, Zfr, 147 anonymous edits
Image Sources, Licenses and Contributors image:Eu-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Eu-TableImage.png License: GNU Free Documentation License Contributors: Paddy, Paginazero, Saperaud Image: EU5P17G-crop.jpg Source: http://en.wikipedia.org/w/index.php?title=File:EU5P17G-crop.jpg License: unknown Contributors: User:Materialscientist File:Eu-Block.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Eu-Block.jpg License: Creative Commons Attribution-Sharealike 3.0 Germany Contributors: user:Alchemist-hp File:Monazit - Mosambik, O-Afrika.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Monazit_-_Mosambik,_O-Afrika.jpg License: unknown Contributors: User:Ra'ike Image:Aperture Grille.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Aperture_Grille.jpg License: Public Domain Contributors: Original uploader was Loongyh at en.wikipedia (Original text : Loongyh (talk))
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
9
Gadolinium
1
Gadolinium 64
europium ← gadolinium → terbium
↑
Gd ↓
Cm Periodic Table - Extended Periodic Table
General Name, Symbol, Number
gadolinium, Gd, 64
Element category
lanthanides
Group, Period, Block
n/a, 6, f
Appearance
silvery white
Standard atomic weight
157.25(3) g·mol
Electron configuration
[Xe] 4f 5d 6s
−1
7
Electrons per shell
1
2
2, 8, 18, 25, 9, 2 Physical properties
Phase
solid
Density (near r.t.)
7.90 g·cm−3
Liquid density at m.p.
7.4 g·cm−3
Melting point
1585 K (1312 °C, 2394 °F)
Boiling point
3546 K (3273 °C, 5923 °F)
Heat of fusion
10.05 kJ·mol−1
Heat of vaporization
301.3 kJ·mol−1
Specific heat capacity
(25 °C) 37.03 J·mol−1·K−1 Vapor pressure (calculated)
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1836
2028
2267
2573
2976
3535
Atomic properties
Gadolinium
2
Crystal structure
hexagonal
Oxidation states
1, 2, 3 (mildly basic oxide)
Electronegativity
1.20 (Pauling scale)
Ionization energies (more)
1st: 593.4 kJ·mol−1 2nd: 1170 kJ·mol−1 3rd: 1990 kJ·mol−1
Atomic radius
180 pm
Covalent radius
196±6 pm Miscellaneous
Magnetic ordering
ferromagnetic/paramagnetic [1] transition at 292 K
Electrical resistivity
(r.t.) (α, poly) 1.310 µΩ·m
Thermal conductivity
(300 K) 10.6 W·m
Thermal expansion
(100 °C) (α, poly) 9.4 µm/(m·K)
Speed of sound (thin rod)
(20 °C) 2680 m/s
Young's modulus
(α form) 54.8 GPa
Shear modulus
(α form) 21.8 GPa
Bulk modulus
(α form) 37.9 GPa
Poisson ratio
−1
·K
−1
(α form) 0.259
Vickers hardness
570 MPa
CAS registry number
7440-54-2 Most-stable isotopes
Main article: Isotopes of gadolinium iso
NA
half-life
152
0.20%
1.08×1014 y
154
2.18%
154
155
14.80%
155
156
20.47%
156
157
15.65%
157
158
24.84%
158
160
21.86%
Gd Gd Gd Gd Gd Gd Gd
DM α
DE (MeV)
DP
2.205
148
1.7
160
Sm
Gd is stable with 90 neutron Gd is stable with 91 neutron Gd is stable with 92 neutron Gd is stable with 93 neutron Gd is stable with 94 neutron
>1.3×1021y
β−β− References
Dy
Gadolinium
3
Gadolinium (pronounced /ˌɡædəˈlɪniəm/) is a chemical element that has the symbol Gd and atomic number 64. It is a silvery-white, malleable and ductile rare-earth metal. Gadolinium has exceptionally high absorption of neutrons and therefore is used for shielding in neutron radiography and in nuclear reactors. Because of its paramagnetic properties, solutions of organic gadolinium complexes and gadolinium compounds are the most popular intravenous MRI contrast agents in medical magnetic resonance imaging.
Characteristics Physical Gadolinium is a silvery-white, malleable and ductile rare-earth metal. It crystallizes in hexagonal, close-packed alpha form at room temperature, but, when heated to temperatures above 1235 °C, it transforms into its beta form, which has a body-centered cubic structure.[2] Gadolinium-157
A yellower sample of gadolinium
has
the
highest
thermal
neutron
capture cross-section of any known nuclide with the exception of xenon-135, 49,000 barns, but it also has a fast burn-out rate, limiting its usefulness as a nuclear control rod material.[3]
Gadolinium is strongly paramagnetic at room temperature, and exhibits ferromagnetic properties below room temperature. Gadolinium demonstrates a magnetocaloric effect whereby its temperature increases when it enters a magnetic field and decreases when it leaves the magnetic field. The effect is considerably stronger for the gadolinium alloy Gd5(Si2Ge2) [4] . Individual gadolinium atoms have been isolated by encapsulating them into fullerene molecules and visualized with transmission electron microscope.[5] . Individual Gd atoms and small Gd clusters have also been incorporated into carbon nanotubes.[6]
Chemical Unlike other rare earth elements, gadolinium is relatively stable in dry air. However, it tarnishes quickly in moist air, forming a loosely-adhering oxide which spalls off, exposing more surface to oxidation. 4 Gd + 3 O2 → 2 Gd2O3 Gadolinium is a strong reducing agent, which reduces oxides of several metals, such as Fe, Cr, Sn, Pb, Mn and Zr, into their elements.[2] Gadolinium is quite electropositive and reacts slowly with cold water and quite quickly with hot water to form gadolinium hydroxide: 2 Gd (s) + 6 H2O (g) → 2 Gd(OH)3 (aq) + 3 H2 (g) Gadolinium metal reacts with all the halogens at temperature about 200 °C: 2 Gd (s) + 3 F2 (g) → 2 GdF3 (s) [white] 2 Gd (s) + 3 Cl2 (g) → 2 GdCl3 (s) [white] 2 Gd (s) + 3 Br2 (g) → 2 GdBr3 (s) [white]
Gadolinium 2 Gd (s) + 3 I2 (g) → 2 GdI3 (s) [yellow] Gadolinium dissolves readily in dilute sulphuric acid to form solutions containing the colorless Gd(III) ions, which exist as a [Gd(OH2)9]3+ complexes:[7] 2 Gd (s) + 3 H2SO4 (aq) → 2 Gd3+ (aq) + 3 SO2−4 (aq) + 3 H2 (g)
Gadolinium combines with nitrogen, carbon, sulfur, phosphorus, boron, selenium, silicon and arsenic at elevated temperatures, forming binary compounds.[2] In those compounds, Gd mostly exhibit oxidation state +3. Gadolinium(II) halogenides are obtained by annealing Gd(III) halogenides in presence of metallic Gd in tantalum containers. Gadolinium also form sesquichloride Gd2Cl3, which can be further reduced to GdCl by annealing at 800 °C. This gadolinium(I) chloride forms platelets with layered graphite-like structure.[8]
Compounds Compounds of gadolinium include: • Fluorides: GdF3 • Chlorides: GdCl3 • Bromides: GdBr3 • • • • • •
Nitrates: Gd(NO3)3 Iodides: GdI3 Oxides: Gd2O3 Sulfides: Gd2S3 Nitrides: GdN Organics: gadodiamide
Isotopes Naturally occurring gadolinium is composed of 6 stable isotopes, 154Gd, 155Gd, 156Gd, 157 Gd, 158Gd and 160Gd, and 1 radioisotope, 152Gd, with 158Gd being the most abundant (24.84% natural abundance). The predicted double beta decay of 160Gd has never been observed (only lower limit on its half-life of more than 1.3×1021 years has been set experimentally [9] ). Twenty nine radioisotopes have been characterized, with the most stable being alpha-decaying 152Gd (naturally occurring) with a half-life of 1.08×1014 years, and 150Gd with a half-life of 1.79×106 years. All of the remaining radioactive isotopes have half-lives less than 74.7 years. The majority of these have half-lives less than 24.6 seconds. Gadolinium isotopes have 4 metastable isomers, with the most stable being 143mGd (T½=110 seconds), 145mGd (T½=85 seconds) and 141mGd (T½=24.5 seconds). The primary decay mode at atomic masses lower than the most abundant stable isotope, 158 Gd, is electron capture, and the primary mode at higher atomic masses is beta decay. The primary decay products for isotopes of weights lower than 158Gd are the element Eu (europium) isotopes and the primary products at higher weights are the element Tb (terbium) isotopes.
4
Gadolinium
5
History In 1880, Swiss chemist Jean Charles Galissard de Marignac observed spectroscopic lines due to gadolinium in samples of didymium and gadolinite; French chemist Paul Émile Lecoq de Boisbaudran separated gadolinia, the oxide of Gadolinium, from Mosander's yttria in 1886. The element itself was isolated only recently. Gadolinium, like the mineral gadolinite, is named after Finnish chemist and geologist Johan Gadolin. In older literature, the natural form of the element is often called an earth, meaning that the element came from Earth.[2]
Occurrence Gadolinium is never found in nature as the free element, but is contained in many rare minerals such as monazite and bastnäsite. It occurs only in trace amounts in the mineral gadolinite, which was also named after Johan Gadolin. The abundance in the earth crust is about 6.2 mg/kg.[2]
Production Gadolinium is produced both from monazite and
Gadolinite
bastnäsite. Crushed minerals are attacked by hydrochloric or sulfuric acid that transforms insoluble rare-earth oxides into soluble chlorides or sulfates. The acidic filtrates are partially neutralized with caustic soda to pH 3-4. Thorium precipitates out of solution as hydroxide and is removed. After that the solution is treated with ammonium oxalate to convert rare earths in to their insoluble oxalates. The oxalates are converted to oxides by annealing. The oxides are dissolved in nitric acid that excludes one of the main components, cerium, whose oxide is insoluble in HNO3. The solution is treated with magnesium nitrate to produce a crystallized mixture of double salts of gadolinium, samarium and europium. The salts are separated by ion exchange. In this process, rare-earth ions are sorbed onto suitable ion-exchange resin by exchange with hydrogen, ammonium or cupric ions present in the resin. The rare earth ions are then selectively washed out by suitable complexing agent.[2] Gadolinium metal is obtained from its oxide or salts by heating with calcium at 1450 °C under argon atmosphere. Sponge gadolinium can be produced by reducing molten GdCl3 with an appropriate metal oxide at temperatures below 1312 °C (melting point of Gd) in a reduced pressure.[2]
Applications Because of extremely high neutron cross-section of gadolinium, this element is very effective for use with neutron radiography and in shielding of nuclear reactors. It is used as a secondary, emergency shut-down measure in some nuclear reactors, particularly of the CANDU type.[2] Gadolinium is also used in nuclear marine propulsion systems as a burnable poison. Gadolinium is an efficient catalyst used for decarboxylation of oxaloacetic acid, convertion of orto- to para-hydrogen and polymerization of ethylene.[2]
Gadolinium Gadolinium also possesses unusual metallurgic properties, with as little as 1% of gadolinium improving the workability and resistance of iron, chromium, and related alloys to high temperatures and oxidation. Because of their paramagnetic properties, solutions of organic gadolinium complexes and gadolinium compounds are used as intravenous MRI contrast agent to enhance images in medical magnetic resonance imaging. Magnevist is the most widespread example.[10] Beside MRI, gadolinium (Gd) is also used in other imaging. In X-ray, gadolinium is contained in the phosphor layer, suspending in a polymer matrix at the detector. Terbium-doped gadolinium oxysulfide (Gd2O2S: Tb) at the phosphor layer is to convert the X-rays releasing from Gadolinium-153 helps calibrate positron emission tomography (PET) systems that are used in nuclear medicine the source into light. This material for functional imaging. This PET image of the human brain emits green light at 540 nm due to the shows the difference between a normal brain and the 3+ presence of Tb , which is very useful clinically depressed patient. The blue color indicates less for enhancing the imaging quality of glucose metabolism in a normal brain. The green, yellow, and red colors indicate areas of higher glucose metabolism the X-ray that is exposed to the [11] characteristic of a depressed patient. photographic film. The energy conversion of Gd is up to 20%, which means, one-fifth of the X-ray striking on the phosphor layer can be converted into light photons. Gadolinium oxyorthosilicate (Gd2SiO5, GSO; usually doped by 0.1-1% of Ce) is a single crystal that is used as a scintillator in medical imaging such as positron emission tomography or for detecting neutrons.[12] Gadolinium-153 is produced in a nuclear reactor from elemental europium or enriched gadolinium targets. It has a half-life of 240±10 days and emits gamma radiation with strong peaks at 41 keV and 102 keV. It is used in many quality assurance applications, such as line sources and calibration phantoms, to ensure that nuclear medicine imaging systems operate correctly and produce useful images of radioisotope distribution inside the patient.[11] It is also used as a gamma ray source in X-ray absorption measurements or in bone density gauges for osteoporosis screening, as well as in the Lixiscope portable X-ray imaging system.[13] Gadolinium is used for making gadolinium yttrium garnet (Gd3Ga5O12); it has microwave applications and is used in fabrication of various optical components and as substrate material for magneto–optical films. Gadolinium compounds are also used for making phosphors for colour TV tubes, compact discs and computer memory.[14]
Biological role Gadolinium has no known native biological role, but in research on biological systems it has a few roles. It is used as a component of MRI contrast agents, as, in the 3+ oxidation state, the metal has 7 unpaired f electrons. This causes water around the contrast agent to relax quickly, enhancing the quality of the MRI scan. Second, as a member of the lanthanides, it is used in various ion channel electrophysiology experiments, where it is used to block sodium leak channels, as well as to stretch activated ion channels.[15]
6
Gadolinium
7
Safety As a free ion, gadolinium is highly toxic but is generally regarded as safe when administered as a chelated compound. The compounds can be classified by whether they are macro-cyclic or linear geometry and whether they are ionic or not. Cyclical ionic Gd compounds being considered the least likely to release the Gd ion and hence the most safe[16] . US Food and Drug Administration approved Gd chelated contrast agents include: Omniscan, Multihance, Magnevist, ProHance, Vasovist and OptiMARK.[17] Gadolinium MRI contrast agents have proved safer than the iodinated contrast agents used in X-ray radiography or computed tomography. Anaphylactoid reactions are rare, occurring in approx. 0.03-0.1%.[18] Although gadolinium agents have proved useful for patients with renal impairment, in patients with severe renal failure requiring dialysis there is a risk of a rare but serious illnesses, such as nephrogenic systemic fibrosis[19] and nephrogenic fibrosing dermopathy[20] , that may be linked to the use of certain gadolinium-containing agents. Although a causal link has not been definitively established, current guidelines in the United States are that dialysis patients should only receive gadolinium agents where essential, and that dialysis should be performed as soon as possible after the scan is complete, in order to remove the agent from the body promptly.[21]
External links • • • • • •
WebElements.com – Gadolinium [22] Nephrogenic Systemic Fibrosis – Complication of Gadolinium MR Contrast [23] MRI Side Effects and Gadolinium Lawsuits [24] It's Elemental – Gadolinium [25] refrigerator uses gadolinium metal that heats up when exposed to magnetic field FDA Advisory on Gadolinium-Based Contrast [27]
[26]
References [1] Charles Kittel (1996). Introduction to Solid State Physics. New York: Wiley. p. 449. [2] Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds (http:/ / books. google. com/ books?id=Xqj-TTzkvTEC& pg=PA243). McGraw-Hill. pp. 302–306. ISBN 0070494398. . Retrieved 2009-06-06. [3] " Gadolinium (http:/ / www. ncnr. nist. gov/ resources/ n-lengths/ elements/ gd. html)". Neutron News (NIST) 3 (3): 29. 1992. . Retrieved 2009-06-06. [4] Karl Gschneidner, Jr. and Kerry Gibson (2001-12-07). " Magnetic refrigerator successfully tested (http:/ / www. external. ameslab. gov/ news/ release/ 01magneticrefrig. htm)". Ames Laboratory. . Retrieved 2006-12-17. [5] Suenaga, Kazu (2003). "Evidence for the Intramolecular Motion of Gd Atoms in a Gd2@C92 Nanopeapod". Nano Letters 3: 1395. doi: 10.1021/nl034621c (http:/ / dx. doi. org/ 10. 1021/ nl034621c). [6] Hashimoto, A; Yorimitsu, H; Ajima, K; Suenaga, K; Isobe, H; Miyawaki, J; Yudasaka, M; Iijima, S; Nakamura, E (Jun 2004). " Selective deposition of a gadolinium(III) cluster in a hole opening of single-wall carbon nanohorn (http:/ / www. pnas. org/ cgi/ pmidlookup?view=long& pmid=15163794)" (Free full text). Proceedings of the National Academy of Sciences of the United States of America 101 (23): 8527–30. doi: 10.1073/pnas.0400596101 (http:/ / dx. doi. org/ 10. 1073/ pnas. 0400596101). ISSN 0027-8424 (http:/ / worldcat. org/ issn/ 0027-8424). PMID 15163794. PMC: 423227 (http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?tool=pmcentrez& artid=423227). . edit (http:/ / en. wikipedia. org/ wiki/ Template:cite_doi/ 10. 1073. 2fpnas. 0400596101) [7] " Chemical reactions of Gadolinium (https:/ / www. webelements. com/ gadolinium/ chemistry. html)". Webelements. . Retrieved 2009-06-06. [8] Cotton (2007). Advanced inorganic chemistry, 6th ed (http:/ / books. google. com/ books?id=U3MWRONWAmMC& pg=PA1128). Wiley-India. p. 1128. ISBN 8126513381. . [9] F. A. Danevich et al., Quest for double beta decay of
160
Gd and Ce isotopes. Nucl. Phys. A 694(2001)375.
Gadolinium [10] Gary Liney (2006). MRI in clinical practice (http:/ / books. google. com/ books?id=xpCffxNrCXYC& pg=PA13). Springer. pp. 13;30. ISBN 184628161X. . [11] " Gadolinium-153 (http:/ / radioisotopes. pnl. gov/ gadolinium. stm)". Pacific Northwest National Laboratory. . Retrieved 2009-06-06. [12]
http:/ / www. webelements. com/ webelements/ elements/ text/ Gd/ index. html http:/ / rad. usuhs. edu/ medpix/ master. php3?mode=slide_sorter& pt_id=10978& quiz=#top http:/ / gadoliniumsideeffects. com/ index-gadolinium. html http:/ / education. jlab. org/ itselemental/ ele064. html http:/ / www. external. ameslab. gov/ news/ release/ 01magneticrefrig. htm http:/ / www. fda. gov/ cder/ drug/ infopage/ gcca/ qa_200705. htm
8
Article Sources and Contributors
Article Sources and Contributors Gadolinium Source: http://en.wikipedia.org/w/index.php?oldid=305362874 Contributors: (jarbarf), 2over0, Ageoflo, Ahoerstemeier, AlimanRuna, Anypodetos, Archimerged, Ariel., Arkuat, B00P, Baccyak4H, Benbest, Blathnaid, BlueEarth, Borislav Dopudja, Brendan Moody, Brian Huffman, Bryan Derksen, CaptainVindaloo, Carnildo, Ceyockey, Chem-awb, Closedmouth, Communisthamster, Conversion script, DSachan, Danielle dk, Darrien, David Latapie, Db099221, Dmn, Drphilharmonic, Duffman, Eddideigel, Edgar181, EinderiheN, Emperorbma, Evil saltine, Expatinsweden, Femto, Fleela, Foobar, Gene Nygaard, GraemeL, GregorB, Grendelkhan, Groyolo, Hallpriest9, Hankwang, Hardy42, Hdt83, Hermitian, Hippiejesus, Icairns, Ideyal, IvanLanin, J.N., Jaan513, Jaraalbe, JazHambo, Jfdwolff, Jj137, Jmknox, Joanjoc, John, Jose77, Julesd, Kaneko ed, Karl-Henner, Kevin Hughes, Kostisl, Kurykh, Kwamikagami, LA2, LeyteWolfer, Looxix, LuigiManiac, Mac addict, Marc Venot, Mashford, Materialscientist, Mav, Minesweeper, Mtodorov 69, Neparis, Nergaal, Nij90, Nimur, Noobeditor, Nsaa, Patstuart, PeepP, Phil Boswell, Plasticup, Plexust, Poolkris, Powernick50, Psyche825, RTC, Remember, Rholton, Rjwilmsi, Robert Foley, Robert K S, Roberta F., Romanm, Saperaud, SauliH, Sbyrnes321, Schneelocke, Selket, Sengkang, Seungfire, Sfuerst, Shaddack, Shinkolobwe, Sionus, Sl, Sleigh, Someone else, Stan J Klimas, Stemonitis, Steve Hart, Stifynsemons, Swpb, Tagishsimon, Tetracube, Tkstark, V1adis1av, Vsmith, Vuo, Walkerma, Watch37264, Weihao.chiu, Whiner01, Wiki alf, Xiggelee, Yekrats, Yyy, Zereshk, 148 anonymous edits
Image Sources, Licenses and Contributors image:Gd-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Gd-TableImage.png License: unknown Contributors: Paddy, Paginazero, Saperaud, Timichal, Will Pittenger Image: Gadolinium-crop.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Gadolinium-crop.jpg License: unknown Contributors: User:JN, User:Materialscientist File:Gadolinium-2.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Gadolinium-2.jpg License: unknown Contributors: User:Jurii File:Gadolinitas.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Gadolinitas.jpg License: GNU Free Documentation License Contributors: Original uploader was WesternDevil at lt.wikipedia File:Gd-153brain.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Gd-153brain.jpg License: Public Domain Contributors: Darrell Fisher, Pacific Northwest National Laboratory (US department of Energy)
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
9
Terbium
1
Terbium 65
gadolinium ← terbium → dysprosium
↑
Tb ↓
Bk Periodic Table - Extended Periodic Table
General Name, Symbol, Number
terbium, Tb, 65
Element category
lanthanides
Group, Period, Block
n/a, 6, f
Appearance
silvery white
Standard atomic weight
158.92535(2) g·mol
Electron configuration
[Xe] 4f 6s
−1
9
Electrons per shell
2
2, 8, 18, 27, 8, 2 Physical properties
Phase
solid
Density (near r.t.)
8.23 g·cm−3
Liquid density at m.p.
7.65 g·cm−3
Melting point
1629 K (1356 °C, 2473 °F)
Boiling point
3503 K (3230 °C, 5846 °F)
Heat of fusion
10.15 kJ·mol−1
Heat of vaporization
293 kJ·mol−1
Specific heat capacity
(25 °C) 28.91 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1789
1979
(2201)
(2505)
(2913)
(3491)
Atomic properties
Terbium
2
Crystal structure
hexagonal
Oxidation states
4, 3, 2, 1 (weakly basic oxide)
Electronegativity
? 1.2 (Pauling scale)
Ionization energies (more)
1st: 565.8 kJ·mol−1 2nd: 1110 kJ·mol−1 3rd: 2114 kJ·mol−1
Atomic radius
177 pm
Covalent radius
194±5 pm Miscellaneous
Magnetic ordering
paramagnetic at 300 K
Electrical resistivity
(r.t.) (α, poly) 1.150 µΩ·m
Thermal conductivity
(300 K) 11.1 W·m
Thermal expansion
(r.t.) (α, poly) 10.3 µm/(m·K)
Speed of sound (thin rod)
(20 °C) 2620 m/s
Young's modulus
(α form) 55.7 GPa
Shear modulus
(α form) 22.1 GPa
Bulk modulus
(α form) 38.7 GPa
Poisson ratio
−1
·K
−1
(α form) 0.261
Vickers hardness
863 MPa
Brinell hardness
677 MPa
CAS registry number
7440-27-9 Most-stable isotopes
Main article: Isotopes of terbium iso
NA
half-life
157
syn
71 y
158
syn
180 y
Tb Tb
159
Tb
100%
DM ε
DE (MeV)
DP
0.060
157
ε
1.220
158
β−
0.937
158
Gd Gd Dy
159
Tb is stable with 94 neutron References
Terbium (pronounced /ˈtɜrbiəm/) is a chemical element with the symbol Tb and atomic number 65. It is a silvery-white rare earth metal that is malleable, ductile and soft enough to be cut with a knife. Terbium is never found in nature as a free element, but it is contained in many minerals, including cerite, gadolinite, monazite, xenotime and euxenite.
Terbium
3
Terbium is used to dope calcium fluoride, calcium tungstate and strontium molybdate, materials that are used in solid-state devices, and as a crystal stabilizer of fuel cells which operate at elevated temperatures. As a component of Terfenol-D (an alloy which expands and contracts in magnetic field more than any other alloy), terbium is of use in actuators, in naval sonar systems and sensors. Terbium oxide is used in green phosphors in fluorescent lamps and color TV tubes. Terbium "green" phosphors (which fluoresce a brilliant lemon-yellow) are combined with divalent Europium blue phosphors and trivalent europium red phosphors to provide the "trichromatic" lighting technology, which is by far the largest consumer of the world's terbium supply.
Characteristics Physical Terbium is a silvery-white rare earth metal that is malleable, ductile and soft enough to be cut with a knife. It is reasonably stable in air (it does not tarnish after nineteen months at room temperature),[1] and two crystal allotropes exist, with a transformation temperature of 1289 °C.[2] The terbium(III) cation is brilliantly fluorescent, in a bright lemon-yellow color that is the result of a strong green emission line in combination with other lines in the orange and red. The yttrofluorite variety of the mineral fluorite owes its creamy-yellow fluorescence in part to terbium. Terbium easily oxidizes and therefore used as element only for research purpose. For example, single Tb atoms have been isolated by implanting them into fullerene molecules.[3] Terbium has a simple ferromagnetic ordering at temperatures below 219 K. Above 219 K, it turns into an helical antiferromagnetic state in which all of the atomic moments in a particular basal plane layer are parallel, and oriented at a fixed angle to the moments of adjacent layers. This unusual antiferromagnetism transforms into a disordered (paramagnetic) state at 230 K.[4]
Chemical The most common valence state of terbium is +3, as in Tb2O3. The +4 state is known in TbO2 and TbF4.[5] [6] Terbium burns readily to form a mixed terbium(III,IV) oxide: 8 Tb + 7 O2 → 2 Tb4O7 In solution, terbium forms only trivalent ions. Terbium is quite electropositive and reacts slowly with cold water and quite quickly with hot water to form terbium hydroxide: 2 Tb (s) + 6 H2O (g) → 2 Tb(OH)3 (aq) + 3H2 (g) Terbium metal reacts with all the halogens: 2 Tb (s) + 3 F2 (g) → 2 TbF3 (s) [white] 2 Tb (s) + 3 Cl2 (g) → 2 TbCl3 (s) [white] 2 Tb (s) + 3 Br2 (g) → 2 TbBr3 (s) [white] 2 Tb (s) + 3 I2 (g) → 2 TbI3 (s) Terbium dissolves readily in dilute sulfuric acid to form solutions containing the pale pink Tb(III) ions, which exist as a [Tb(OH2)9]3+ complexes:[7] 2 Tb (s) + 3 H2SO4 (aq) → 2 Tb3+ (aq) + 3 SO2−4 (aq) + 3 H2 (g)
Terbium
Compounds Terbium combines with nitrogen, carbon, sulfur, phosphorus, boron, selenium, silicon and arsenic at elevated temperatures, forming various binary compounds such as TbH2, TBH3, TbB2, Tb2S3, TbSe, TbTe and TbN.[6] In those compounds, Tb mostly exhibit oxidation states +3 and sometimes +2. Terbium(II) halogenides are obtained by annealing Tb(III) halogenides in presence of metallic Tb in tantalum containers. Terbium also form sesquichloride Tb2Cl3, which can be further reduced to TbCl by annealing at 800 °C. This terbium(I) chloride forms platelets with layered graphite-like structure.[8] Other compounds include • • • •
Chlorides: TbCl3 Bromides: TbBr3 Iodides: TbI3 Fluorides: TbF3, TbF4
Terbium(IV) fluoride is a strong fluorinating agent, emitting relatively pure atomic fluorine when heated[9] rather than the mixture of fluoride vapors emitted from CoF3 or CeF4.
Isotopes Naturally occurring terbium is composed of 1 stable isotope, 159Tb. 33 radioisotopes have been characterized, with the most stable being 158Tb with a half-life of 180 years, 157Tb with a half-life of 71 years, and 160Tb with a half-life of 72.3 days. All of the remaining radioactive isotopes have half-lives that are less than 6.907 days, and the majority of these have half-lives that are less than 24 seconds. This element also has 18 meta states, with the most stable being 156m1Tb (t½ 24.4 hours), 154m2Tb (t½ 22.7 hours) and 154m1Tb (t½ 9.4 hours). The primary decay mode before the most abundant stable isotope, 159Tb, is electron capture, and the primary mode after is beta minus decay. The primary decay products before 159Tb are element Gd (gadolinium) isotopes, and the primary products after are element Dy (dysprosium) isotopes.
History Terbium was discovered in 1843 by Swedish chemist Carl Gustaf Mosander, who detected it as an impurity in Yttrium oxide, Y2O3, and named after the village Ytterby in Sweden. It was not isolated in pure form until the recent advent of ion exchange techniques.[10] When Mosander first partitioned "yttria" into three fractions, "terbia" was the fraction that contained the pink color (due to what is now known as erbium), and "erbia" was the fraction that was essentially colorless in solution, but gave a brown-tinged oxide. Later workers had difficulty in observing the latter, but the pink fraction was impossible to miss. Arguments went back and forth as to whether "erbia" even existed. In the confusion, the original names got reversed, and the exchange of names stuck. It is now thought that those workers who used the double sodium or potassium sulfates to remove "ceria" from "yttria" inadvertently lost the terbium content of the system into the ceria-containing precipitate. In any case, what is now known as terbium was only about 1% of the original yttria, but that was sufficient to impart a yellowish color to the oxide. Thus, terbium was a minor component in the original terbium fraction, dominated by its immediate neighbors, gadolinium and dysprosium. Thereafter, whenever other rare earths were teased apart from
4
Terbium this mixture, whichever fraction gave the brown oxide retained the terbium name, until at last it was pure. The 19th century investigators did not have the benefit of fluorescence technology, wherewith to observe the brilliant fluorescence that would have made this element much easier to track in mixtures.[10]
Occurrence Terbium is never found in nature as a free element, but it is contained in many minerals, including cerite, gadolinite, monazite ((Ce,La,Th,Nd,Y)PO4, which contains up to 0.03% of terbium), xenotime (YPO4) and euxenite ((Y,Ca,Er,La,Ce,U,Th)(Nb,Ta,Ti)2O6, which contains 1% or more of terbium). The crust abundance of terbium is estimated as 1.2 mg/kg.[6] The richest current commercial sources of terbium are the ion-adsorption clays of southern China. The high-yttrium concentrate versions of these are about Xenotime two-thirds yttrium oxide by weight, and about 1% terbia. However, small amounts occur in bastnäsite and monazite, and when these are processed by solvent-extraction to recover the valuable heavy lanthanides in the form of "samarium-europium-gadolinium concentrate" (SEG concentrate), the terbium content of the ore ends up therein. Due to the large volumes of bastnäsite processed, relative to the richer ion-adsorption clays, a significant proportion of the world's terbium supply comes from bastnäsite.[2]
Production Crushed terbium-containing minerals are treated with hot concentrated sulfuric acid to produce water-soluble sulfates of rare earths. The acidic filtrates are partially neutralized with caustic soda to pH 3-4. Thorium precipitates out of solution as hydroxide and is removed. After that the solution is treated with ammonium oxalate to convert rare earths in to their insoluble oxalates. The oxalates are converted to oxides by annealing. The oxides are dissolved in nitric acid that excludes one of the main components, cerium, whose oxide is insoluble in HNO3. Terbium is separated as a double salt with ammonium nitrate by crystallization. [6] The most efficient separation routine for terbium salt from the rare-earth salt solution is ion exchange. In this process, rare-earth ions are sorbed onto suitable ion-exchange resin by exchange with hydrogen, ammonium or cupric ions present in the resin. The rare earth ions are then selectively washed out by suitable complexing agent. As with other rare earths, terbium metal is produced by reducing the anhydrous chloride or fluoride with calcium metal. Calcium and tantalum impurities can be removed by vacuum remelting, distillation, amalgam formation or zone melting.[6]
5
Terbium
6
Applications Terbium is used to dope calcium fluoride, calcium tungstate and strontium molybdate, materials that are used in solid-state devices, and as a crystal stabilizer of fuel cells which operate at elevated temperatures, together with ZrO2.[2] Terbium is also used in alloys and in the production of electronic devices. As a component of Terfenol-D, terbium is of use in actuators, in naval sonar systems, sensors, in the SoundBug device (its first commercial application), and other magnetomechanical devices. Terfenol-D is an alloy that expands or contracts in the presence of a magnetic field. It has the highest magnetostriction of any alloy.[11] Terbium oxide is used in green phosphors in fluorescent lamps and color TV tubes. Sodium terbium borate is used in solid state devices. The brilliant fluorescence allows terbium to be used as a probe in biochemistry, where it somewhat resembles calcium in its behavior. Terbium "green" phosphors (which fluoresce a brilliant lemon-yellow) are combined with divalent Europium blue phosphors and trivalent europium red phosphors to provide the "trichromatic" lighting technology, which is by far the largest consumer of the world's terbium supply. Trichromatic lighting provides much higher light output for a given amount of electrical energy than does incandescent lighting.[2]
Precautions As with the other lanthanides, terbium compounds are of low to moderate toxicity, although their toxicity has not been investigated in detail. Terbium has no known biological role.[2]
See also External links • WebElements.com – Terbium • It's Elemental – Terbium [13]
[12]
pnb:میبرٹ
References [1] " Rare-Earth Metal Long Term Air Exposure Test (http:/ / www. elementsales. com/ re_exp/ index. htm)". . Retrieved 2009-05-05. [2] C. R. Hammond, "The Elements", in Handbook of Chemistry and Physics 81th edition, CRC press. [3] Shimada, T (2004). "Transport properties of C78, C90 and Dy@C82 fullerenes-nanopeapods by field effect transistors". Physica E Low-dimensional Systems and Nanostructures 21: 1089. doi: 10.1016/j.physe.2003.11.197 (http:/ / dx. doi. org/ 10. 1016/ j. physe. 2003. 11. 197). edit (http:/ / en. wikipedia. org/ wiki/ Template:cite_doi/ 10. 1016. 2fj. physe. 2003. 11. 197) [4] M. Jackson (2000). " Magnetism of Rare Earth (http:/ / www. irm. umn. edu/ quarterly/ irmq10-3. pdf)". The IRM quarterly 10 (3): 1. . [5] D.M. Gruen, W.C. Koehler, and J.J. Katz (April 1951). " Higher Oxides of the Lanthanide Elements: Terbium Dioxide (http:/ / pubs. acs. org/ cgi-bin/ abstract. cgi/ jacsat/ 1951/ 73/ i04/ f-pdf/ f_ja01148a020. pdf)" (PDF). Journal of the American Chemical Society: 1475. . [6] Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds (http:/ / books. google. com/ books?id=Xqj-TTzkvTEC& pg=PA243). McGraw-Hill. pp. 920–921. ISBN 0070494398. . Retrieved 2009-06-06. [7] " Chemical reactions of Terbium (https:/ / www. webelements. com/ terbium/ chemistry. html)". Webelements. . Retrieved 2009-06-06. [8] Cotton (2007). Advanced inorganic chemistry, 6th ed (http:/ / books. google. com/ books?id=U3MWRONWAmMC& pg=PA1128). Wiley-India. p. 1128. ISBN 8126513381. .
Terbium [9] J.V.Rau; N.S. Chilingarov, M.S. Leskiv, V.F. Sukhoverkhov', V. Rossi Albertini, L.N. Sidorov (2001). Transition and rare earth metal fluorides as thermal sources of atomic and molecular fluorine. [10] C. K. Gupta, Nagaiyar Krishnamurthy (2004). Extractive metallurgy of rare earths (http:/ / books. google. com/ books?id=E1npz8pwmYwC& pg=PA5). CRC Press. p. 5. ISBN 0415333407. . [11] Rodriguez, C (2009). "New elastomer–Terfenol-D magnetostrictive composites". Sensors and Actuators a Physical 149: 251. doi: 10.1016/j.sna.2008.11.026 (http:/ / dx. doi. org/ 10. 1016/ j. sna. 2008. 11. 026). [12] http:/ / www. webelements. com/ webelements/ elements/ text/ Tb/ index. html [13] http:/ / education. jlab. org/ itselemental/ ele065. html
7
Article Sources and Contributors
Article Sources and Contributors Terbium Source: http://en.wikipedia.org/w/index.php?oldid=308846010 Contributors: 2help, Addshore, Ahoerstemeier, AlimanRuna, Anastrophe, Animum, Archimerged, Arkuat, Benbest, BjKa, BlueEarth, Bobblewik, Borislav Dopudja, Bryan Derksen, Bubba hotep, Cacycle, Carnildo, Conversion script, Crystal whacker, Darrien, David Latapie, Dead3y3, DocKrin, Donarreiskoffer, Dycedarg, DÅ‚ugosz, Edgar181, Emperorbma, Evil saltine, Fastily, FelisSchrödingeris, Femto, Fivemack, Grendelkhan, Grunt, Hankwang, Helge Skjeveland, Heron, Icairns, InfoCan, Iridescent, IvanLanin, Jaan513, Jaraalbe, Joanjoc, Jose77, Karl-Henner, Keilana, Kelovy, Korath, Kurykh, Kwamikagami, LA2, Laurinavicius, Lexicon, Looxix, Marc Venot, Materialscientist, Mav, Minesweeper, Mistercow, NCurse, Nergaal, Nlu, PeepP, Plexust, Polyparadigm, Poolkris, Pras, RTC, Remember, Roberta F., Romanm, Ruakh, Saperaud, Schneelocke, Sl, Squids and Chips, Steve Hart, Stifynsemons, Suisui, Sunnyoraish, Svante, Tagishsimon, Tetracube, Toon05, Uncle G, Vicki Rosenzweig, Vsmith, Walkerma, Warut, WillMak050389, Yekrats, Yyy, 66 anonymous edits
Image Sources, Licenses and Contributors image:Tb-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Tb-TableImage.png License: GNU Free Documentation License Contributors: user:Schneelocke Image: Terbium-croprotated.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Terbium-croprotated.jpg License: unknown Contributors: User:Materialscientist Image:Xenotim mineralogisches museum bonn.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Xenotim_mineralogisches_museum_bonn.jpg License: unknown Contributors: User:Elya
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
8
Dysprosium
1
Dysprosium 66
terbium ← dysprosium → holmium
↑
Dy ↓
Cf Periodic Table - Extended Periodic Table
General Name, Symbol, Number
dysprosium, Dy, 66
Element category
lanthanides
Group, Period, Block
n/a, 6, f
Appearance
silvery white
Standard atomic weight
162.500(1) g·mol
Electron configuration
[Xe] 4f
−1
10
Electrons per shell
2
6s
2, 8, 18, 28, 8, 2 Physical properties
Phase
solid
Density (near r.t.)
8.540 g·cm−3
Liquid density at m.p.
8.37 g·cm−3
Melting point
1680 K (1407 °C, 2565 °F)
Boiling point
2840 K (2562 °C, 4653 °F)
Heat of fusion
11.06 kJ·mol−1
Heat of vaporization
280 kJ·mol−1
Specific heat capacity
(25 °C) 27.7 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1378
1523
(1704)
(1954)
(2304)
(2831)
Atomic properties
Dysprosium
2
Crystal structure
hexagonal
Oxidation states
3, 2 (weakly basic oxide)
Electronegativity
1.22 (Pauling scale)
Ionization energies (more)
1st: 573.0 kJ·mol−1 2nd: 1130 kJ·mol−1 3rd: 2200 kJ·mol−1
Atomic radius
178 pm
Covalent radius
192±7 pm Miscellaneous
Magnetic ordering
paramagnetic at 300 K
Electrical resistivity
(r.t.) (α, poly) 926 nΩ·m
Thermal conductivity
(300 K) 10.7 W·m
Thermal expansion
(r.t.) (α, poly) 9.9 µm/(m·K)
Speed of sound (thin rod)
(20 °C) 2710 m/s
Young's modulus
(α form) 61.4 GPa
Shear modulus
(α form) 24.7 GPa
Bulk modulus
(α form) 40.5 GPa
Poisson ratio
−1
·K
−1
(α form) 0.247
Vickers hardness
540 MPa
Brinell hardness
500 MPa
CAS registry number
7429-91-6 Most-stable isotopes
Main article: Isotopes of dysprosium iso 154
Dy
NA syn
half-life
2.947
150
α
?
152
1×1018y
158
0.10%
158
160
2.34%
160
161
18.91%
161
162
25.51%
162
163
24.90%
163
164
28.18%
164
Dy Dy Dy Dy Dy
DP
α
0.06%
Dy
DE (MeV)
3.0×106y
156
Dy
DM
Dy is stable with 92 neutron Dy is stable with 94 neutron Dy is stable with 95 neutron Dy is stable with 96 neutron Dy is stable with 97 neutron Dy is stable with 98 neutron References
Gd Gd
Dysprosium
3
Dysprosium (pronounced /dɪsˈproʊziəm/) is a chemical element with the symbol Dy and atomic number 66. It is a rare earth element with a metallic silver luster. Dysprosium is never found in nature as a free element, though it is found in various minerals, such as xenotime. Naturally occurring dysprosium is composed of 7 isotopes, the most abundant of 164 which is Dy. Dysprosium was first identified in 1886 by Paul Émile Lecoq de Boisbaudran, but was not isolated in pure form until the development of ion exchange techniques in the 1950s. Dysprosium is used for its high thermal neutron absorption cross-section in making control rods in nuclear reactors, for its high magnetic susceptibility to magnetization in data storage devices and as a component of Terfenol-D. Soluble dysprosium salts are mildly toxic, while the insoluble salts are considered non-toxic.
Characteristics Physical Dysprosium is a rare earth element that has a metallic, bright silver luster. It is soft enough to be cut with a knife, and can be machined without sparking if overheating is avoided. Dysprosium's physical characteristics can be greatly affected even by small amounts of impurities. [1] Dysprosium easily oxidizes and therefore used as element only for research purpose. For example, single Dy atoms have been isolated by implanting them into fullerene molecules.[2] Dysprosium and holmium have the highest magnetic strengths of the elements,[3] especially at low temperatures.[4] Dysprosium has a simple ferromagnetic ordering at temperatures below 85 K. Above 85 K, it turns into an helical antiferromagnetic state in which all of the atomic moments in a particular basal plane layer are parallel, and oriented at a fixed angle to the moments of adjacent layers. This unusual antiferromagnetism transforms into a disordered (paramagnetic) state at 179 K.[5] Dysprosium sample
Chemical Dysprosium metal tarnishes slowly in air and burns readily to form dysprosium(III) oxide: 4 Dy + 3 O2 → 2 Dy2O3 Dysprosium is quite electropositive and reacts slowly with cold water and quite quickly with hot water to form dysprosium hydroxide: 2 Dy (s) + 6 H2O (g) → 2 Dy(OH)3 (aq) + 3 H2 (g) Dysprosium metal vigorously reacts with all the halogens at above 200 °C: 2 Dy (s) + 3 F2 (g) → 2 DyF3 (s) [green] 2 Dy (s) + 3 Cl2 (g) → 2 DyCl3 (s) [white] 2 Dy (s) + 3 Br2 (g) → 2 DyBr3 (s) [white] 2 Dy (s) + 3 I2 (g) → 2 DyI3 (s) [green] Dysprosium dissolves readily in dilute sulfuric acid to form solutions containing the yellow Dy(III) ions, which exist as a [Dy(OH2)9]3+ complexes:[6] 2 Dy (s) + 3 H2SO4 (aq) → 2 Dy3+ (aq) + 3 SO2−4 (aq) + 3 H2 (g)
Dysprosium
Compounds Dysprosium halides, such as DyF3 and DyBr3, tend to take on a yellow color. Dysprosium oxide, also known as dysprosia, is a white powder that is highly magnetic, more so than iron oxide.[4] Dysprosium combines with various non-metals at high temperatures to form binary compounds with varying composition and oxidation states +3 and sometimes +2, such as DyN, DyP, DyH2 and DyH3; DyS, DyS2, Dy2S3 and Dy5S7; DyB2, DyB4, DyB6 and DyB12, as well as Dy3C and Dy2C3.[7] Dysprosium carbonate, Dy2(CO3)3, and dysprosium sulfate, Dy2(SO4)3, result from similar reactions.[8] Most dysprosium compounds are soluble in water, though dysprosium carbonate tetrahydrate (Dy2(CO3)3•4H2O) and dysprosium oxalate decahydrate (Dy2(C2O4)3•10H2O) are both insoluble in water.[9] [10]
Isotopes Naturally occurring dysprosium is composed of 7 isotopes: 156Dy, 158Dy, 160Dy, 161Dy, 162 Dy, 163Dy, and 164Dy. These are all considered stable, although 156Dy decays by alpha decay with a half-life of over 1×1018 years. Of the naturally occurring isotopes, 164Dy is the most abundant at 28%, followed by 162Dy at 26%. The least abundant is 156Dy at .06%.[11] Twenty-nine radioisotopes have also been synthesized, ranging in atomic mass from 138 to 173. The most stable of these is 154Dy with a half-life of approximately 3 × 106 years, followed by 159Dy with a half-life of 144.4 days. The least stable is 138Dy with a half-life of 200 ms. Isotopes that are lighter than the stable isotopes tend to decay primarily by β+ decay, while those that are heavier tend to decay by β- decay, with some exceptions. 154Dy decays primarily by alpha decay, and 152Dy and 159Dy decay primarily by electron capture.[11] Dysprosium also has at least 11 metastable isomers, ranging in atomic mass from 140 to 165. The most stable of these is 165mDy, which has a half-life of 1.257 minutes. 149 Dy has two excitation states, the second of which, 149m2Dy, has a half-life of 28 ns.[11]
History In 1878, erbium ores were found to contain the oxides of two other rare earths: holmium and thulium. French chemist Paul Émile Lecoq de Boisbaudran, while working with holmium oxide, separated dysprosium oxide from it in Paris in 1886.[12] His procedure for isolating the dysprosium involved dissolving dysprosium oxide in acid, then adding ammonia to precipitate the hydroxide. He was only able to isolate dysprosium from its oxide after more than 30 attempts at his procedure. Upon succeeding, he named the element dysprosium from the Greek dysprositos, meaning "hard to get". However, the element was not isolated in relatively pure form until after the development of ion exchange techniques by Frank Spedding at Iowa State University in the early 1950s.[3] In 1950, Glenn T. Seaborg, Albert Ghiorso, and Stanley G. Thompson bombarded 241Am with helium ions, which produced atoms with an atomic number of 97 and which closely resembled the neighboring lanthanide terbium. Because terbium was named after Ytterby, the city in which it and several other elements were discovered, this new element was named berkelium for the city in which it was synthesized. However, when the research team synthesized element 98, they could not think of a good analogy for dysprosium, and instead named the element californium in honor of the state in which it was synthesized.
4
Dysprosium
5
The research team went on to "point out that, in recognition of the fact that dysprosium is named on the basis of a Greek word meaning 'difficult to get at,' that the searchers for another element a century ago found it difficult to get to California."[13]
Occurrence Dysprosium is never encountered as a free element, but is found in many minerals, including xenotime, fergusonite, gadolinite, euxenite, polycrase, blomstrandine, monazite and bastnäsite; often with erbium and holmium or other rare earth elements. Currently, most dysprosium is being obtained from the ion-adsorption clay ores of southern China. In the high-yttrium version of these, dysprosium happens to be the most abundant of the heavy lanthanides, comprising up to 7–8% of the concentrate (as compared to about 65% for yttrium).[14] [15] The concentration of Dy in the Earth crust is about 5.2 mg/kg and in sea water 0.9 ng/L.[7]
Xenotime
Production Dysprosium is obtained primarily from monazite sand, a mixture of various phosphates. The metal is obtained as a by-product in the commercial extraction of yttrium. In isolating dysprosium, most of the unwanted metals can be removed magnetically or by a flotation process. Dysprosium can then be separated from other rare earth metals by an ion exchange displacement process. The resulting dysprosium ions can then react with either fluorine or chlorine to form dysprosium fluoride, DyF3, or dysprosium chloride, DyCl3. These compounds can be reduced using either calcium or lithium metals in the following reactions:[8] 3 Ca + 2 DyF3 → 2 Dy + 3 CaF2 3 Li + DyCl3 → Dy + 3 LiCl The components are placed in a tantalum crucible and fired in a helium atmosphere. As the reaction progresses, the resulting halide compounds and molten dysprosium separate due to differences in density. When the mixture cools, the dysprosium can be cut away from the impurities.[8] About 100 tonnes of dysprosium are produced worldwide each year.[16]
Applications Dysprosium is used, in conjunction with vanadium and other elements, in making laser materials. Because of dysprosium's high thermal neutron absorption cross-section, dysprosium oxide-nickel cermets are used in neutron-absorbing control rods in nuclear reactors.[17] Dysprosium-cadmium chalcogenides are sources of infrared radiation which is useful for studying chemical reactions.[1] Because dysprosium and its compounds are highly susceptible to magnetization, they are employed in various data storage applications, such as in compact discs.[18]
Dysprosium Neodymium-iron-boron magnets can have up to 6% of the neodymium substituted with dysprosium[19] to raise the coercivity for demanding applications such as drive motors for hybrid electric vehicles. This substitution would require up to 100 grams of dysprosium per hybrid car produced. Based on Toyota's projected 2 million units per year, the use of dysprosium in applications such as this would quickly exhaust the available supply of the metal.[20] The dysprosium substitution may also be useful in other applications, as it improves the corrosion resistance of the magnets.[21] Dysprosium is one of the components of Terfenol-D, along with iron and terbium. Terfenol-D has the highest room-temperature magnetoresistance of any known material;[22] this property is employed in transducers, wide-band mechanical resonators,[23] and high-precision liquid fuel injectors.[24] Dysprosium is used in dosimeters for measuring ionizing radiation. Crystals of calcium sulfate or calcium fluoride are doped with dysprosium. When these crystals are exposed to radiation, the dysprosium atoms become excited and luminescent. The luminescence can be measured to determine the degree of exposure to which the dosimeter has been subjected.[3] Nanofibers of dysprosium compounds have high Nanofibers of dysprosium oxide fluoride strength and large surface area; therefore, they can be used for reinforcement of other materials and as a catalyst. Fibers of dysprosium oxide fluoride can be produced by heating an aqueous solution of DyBr and NaF to 450 °C at 450 bar pressure for 17 hours. This material is remarkably robust, surviving over 100 hours in various aqueous solutions at temperatures exceeding 400 °C without re-dissolving or aggregating.[25] [26] [27]
Precautions Like many powders, dysprosium powder may present an explosion hazard when mixed with air and when an ignition source is present. Thin foils of the substance can also be ignited by sparks or by static electricity. Dysprosium fires cannot be put out by water. It can react with water to produce flammable hydrogen gas.[28] Dysprosium chloride fires, however, can be extinguished with water,[29] while dysprosium fluoride and dysprosium oxide are non-flammable.[30] [31] Dysprosium nitrate, Dy(NO3)3, is a strong oxidizing agent and will readily ignite upon contact with organic substances.[4] Soluble dysprosium salts, such as dysprosium chloride and dysprosium nitrate, are mildly toxic when ingested. The insoluble salts, however, are non-toxic. Based on the toxicity of dysprosium chloride to mice, it is estimated that the ingestion of 500 grams or more could be fatal to a human.[3]
6
Dysprosium
7
See also • Lanthanoid
External links • WebElements.com – Dysprosium • It's Elemental – Dysprosium [33]
[32]
References [1] Lide, David R., ed (2007–2008). "Dysprosium". CRC Handbook of Chemistry and Physics. 4. New York: CRC Press. pp. 11. ISBN 978-0-8493-0488-0. [2] Shimada, T (2004). "Transport properties of C78, C90 and Dy@C82 fullerenes-nanopeapods by field effect transistors". Physica E Low-dimensional Systems and Nanostructures 21: 1089. doi: 10.1016/j.physe.2003.11.197 (http:/ / dx. doi. org/ 10. 1016/ j. physe. 2003. 11. 197). edit (http:/ / en. wikipedia. org/ wiki/ Template:cite_doi/ 10. 1016. 2fj. physe. 2003. 11. 197) [3] Emsley, John (2001). Nature's Building Blocks. Oxford: Oxford University Press. pp. 129–132. ISBN 0-19-850341-5. [4] Krebs, Robert E. (1998). "Dysprosium". The History and Use of our Earth's Chemical Elements. Greenwood Press. pp. 234–235. ISBN 0-313-30123-9. [5] Jackson, Mike (2000). (PDF)IRM Quarterly (Institute for Rock Magnetism) 10 (3): 6. http:/ / www. irm. umn. edu/ quarterly/ irmq10-3. pdf. [6] " Chemical reactions of Dysprosium (https:/ / www. webelements. com/ dysprosium/ chemistry. html)". Webelements. . Retrieved 2009-06-06. [7] Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds (http:/ / books. google. com/ books?id=Xqj-TTzkvTEC& pg=PA243). McGraw-Hill. pp. 289–290. ISBN 0070494398. . Retrieved 2009-06-06. [8] Heiserman, David L. (1992). Exploring Chemical Elements and their Compounds. TAB Books. pp. 236–238. ISBN 0-8306-3018-X. [9] Perry, D. L. (1995). Handbook of Inorganic Compounds. CRC Press. pp. 152–154. ISBN 0-8492-8671-3. [10] Jantsch, G.; Ohl, A. (1911). "Zur Kenntnis der Verbindungen des Dysprosiums". Berichte der deutschen chemischen Gesellschaft 44 (2): 1274–1280. doi: 10.1002/cber.19110440215 (http:/ / dx. doi. org/ 10. 1002/ cber. 19110440215). [11] Audi, G. (2003). "Nubase2003 Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3–128. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [12] Paul Émile Lecoq de Boisbaudran (1886). " L'holmine (ou terre X de M Soret) contient au moins deux radicaux métallique (Holminia contains at least two metal) (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k3058f/ f1001. chemindefer)" (in French). Comptes Rendus 143: 1003–1006. . [13] Weeks, M. E. (1968). Discovery of the Elements (7 ed.). Journal of Chemical Education. pp. 848–849. ISBN 0848685792. OCLC 23991202 (http:/ / worldcat. org/ oclc/ 23991202). [14] Naumov, A. V. (2008). " Review of the World Market of Rare-Earth Metals (http:/ / www. springerlink. com/ content/ y8925j378w4u4175/ )". Russian Journal of Non-Ferrous Metals 49 (1): 14–22. . [15] Gupta, C. K.; Krishnamurthy N. (2005). Extractive Metallurgy of Rare Earths (http:/ / books. google. com/ books?id=F0Bte_XhzoAC). CRC Press. ISBN 9780415333405. . [16] " Dysprosium (Dy) - Chemical properties, Health and Environmental effects (http:/ / www. lenntech. com/ Periodic-chart-elements/ Dy-en. htm)". Lenntech Water treatment & air purification Holding B.V.. 2008. . Retrieved 2009-06-02. [17] Amit, Sinha; Beant Prakash, Sharma (2005). "Development of Dysprosium Titanate Based Ceramics". Journal of the American Ceramic Society 88 (4): 1064–1066. doi: 10.1111/j.1551-2916.2005.00211.x (http:/ / dx. doi. org/ 10. 1111/ j. 1551-2916. 2005. 00211. x). [18] Lagowski, J. J., ed (2004). Chemistry Foundations and Applications. 2. Thomson Gale. pp. 267–268. ISBN 0-02-865724-1. [19] Shi, Fang, X.; Jiles, Y. (1998). "Modeling of magnetic properties of heat treated Dy-doped NdFeBparticles bonded in isotropic and anisotropic arrangements". IEEE Transactions on Magnetics 34 (4): 1291–1293. doi: 10.1109/20.706525 (http:/ / dx. doi. org/ 10. 1109/ 20. 706525). [20] Campbell, Peter (February 2008). " Supply and Demand, Part 2 (http:/ / www. magnetweb. com/ Col05. htm)". Princeton Electro-Technology, Inc.. .
Dysprosium [21] Yu, L. Q.; Wen, Y. H.; Yan, M. (2004). "Effects of Dy and Nb on the magnetic properties and corrosion resistance of sintered NdFeB". Journal of Magnetism and Magnetic Materials 283 (2-3): 353–356. doi: 10.1016/j.jmmm.2004.06.006 (http:/ / dx. doi. org/ 10. 1016/ j. jmmm. 2004. 06. 006). [22] " What is Terfenol-D? (http:/ / etrema-usa. com/ core/ terfenold/ )". ETREMA Products, Inc.. 2003. . Retrieved 2008-11-06. [23] Kellogg, Rick; Flatau, Alison (May 2004). "Wide Band Tunable Mechanical Resonator Employing the ΔE Effect of Terfenol-D". Journal of Intelligent Material Systems & Structures (Sage Publications, Ltd) 15 (5): 355–368. doi: 10.1177/1045389X04040649 (http:/ / dx. doi. org/ 10. 1177/ 1045389X04040649). [24] Leavitt, Wendy (February 2000). " Take Terfenol-D and call me (http:/ / 0-search. ebscohost. com. ilsprod. lib. neu. edu/ login. aspx?direct=true& db=buh& AN=2869368& site=ehost-live)". Fleet Owner (RODI Power Systems Inc) 95 (2): 97. . Retrieved 2008-11-06. [25] " Supercritical Water Oxidation/Synthesis (http:/ / www. pnl. gov/ supercriticalfluid/ tech_oxidation. stm)". Pacific Northwest National Laboratory. . Retrieved 2009-06-06. [26] " Rare Earth Oxide Fluoride: Ceramic Nano-particles via a Hydrothermal Method (http:/ / availabletechnologies. pnl. gov/ technology. asp?id=152)". Pacific Northwest National Laboratory. . Retrieved 2009-06-06. [27] M.M. Hoffman, J.S. Young, J.L. Fulton (2000). "Unusual dysprosium ceramic nano-fiber growth in a supercritical aqueous solution". J Mat. Sci. 35: 4177. doi: 10.1023/A:1004875413406 (http:/ / dx. doi. org/ 10. 1023/ A:1004875413406). [28] Dierks, Steve (January 2003). " Dysprosium (http:/ / www. espi-metals. com/ msds's/ Dysprosium. htm)". Material Safety Data Sheets. Electronic Space Products International. . Retrieved 2008-10-20. [29] Dierks, Steve (January 1995). " Dysprosium Chloride (http:/ / www. espi-metals. com/ msds's/ Dysprosium Chloride. htm)". Material Safety Data Sheets. Electronic Space Products International. . Retrieved 2008-11-07. [30] Dierks, Steve (December 1995). " Dysprosium Fluoride (http:/ / www. espi-metals. com/ msds's/ Dysprosium Fluoride. htm)". Material Safety Data Sheets. Electronic Space Products International. . Retrieved 2008-11-07. [31] Dierks, Steve (November 1988). " Dysprosium Oxide (http:/ / www. espi-metals. com/ msds's/ Dysprosium Oxide. htm)". Material Safety Data Sheets. Electronic Space Products International. . Retrieved 2008-11-07. [32] http:/ / www. webelements. com/ webelements/ elements/ text/ Dy/ key. html [33] http:/ / education. jlab. org/ itselemental/ ele066. html
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Article Sources and Contributors
Article Sources and Contributors Dysprosium Source: http://en.wikipedia.org/w/index.php?oldid=308883418 Contributors: 2over0, AWeishaupt, Ahoerstemeier, AlimanRuna, AndonicO, Andre Engels, Archimerged, Arkuat, Atjesse, Beetstra, Benbest, Betacommand, Bkell, BlueEarth, Borislav Dopudja, BrotherFlounder, Bryan Derksen, Burntsauce, Carnildo, Chasingsol, Chem-awb, Conversion script, Corbon, Corpx, Cryptic C62, Crystal whacker, Darrien, David Latapie, Decadencecavy, Donarreiskoffer, Doonhamer, Dysprosia, DÅ‚ugosz, Edgar181, El C, Emperorbma, Everyguy, Evil saltine, Femto, Fivemack, Grendelkhan, Gurch, Hankwang, Helge Skjeveland, Herbee, Icairns, Ivan05, IvanLanin, J.delanoy, Jaan513, Jaraalbe, Joanjoc, Jose77, Karl-Henner, Kelovy, KnightLago, Kurykh, Kwamikagami, LA2, LarryMorseDCOhio, Lloydpick, Looxix, Mac, Marc Venot, Materialscientist, Mav, Mbruck, Melchoir, Michael Devore, Minesweeper, Nergaal, Nihiltres, Nimur, PROCodyHASKA, Phyzome, Plexust, Poolkris, Possum, Puchiko, RTC, Remember, Rich Farmbrough, Richard Harvey, Roberta F., Romanm, Sanbeg, Saperaud, Schneelocke, Sengkang, Sfuerst, Shadowdemon0085, Sl, Snowolf, Steve Hart, Stifynsemons, Stone, Suisui, Swaq, Tagishsimon, Tbone, Tetracube, Thedemonhog, Theseeker4, Tiptoety, V8rik, Vsmith, Walkerma, Warut, Yekrats, Yyy, 114 anonymous edits
Image Sources, Licenses and Contributors image:Dy-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Dy-TableImage.png License: GNU Free Documentation License Contributors: user:Schneelocke Image: Dy chips.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Dy_chips.jpg License: unknown Contributors: User:Materialscientist Image:Dysprosium.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Dysprosium.jpg License: GNU Free Documentation License Contributors: Original uploader was Tomihahndorf at de.wikipedia Image:Xenotímio1.jpeg Source: http://en.wikipedia.org/w/index.php?title=File:Xenotímio1.jpeg License: Creative Commons Attribution-Sharealike 2.0 Contributors: Zimbres Image:DyFibers2.jpg Source: http://en.wikipedia.org/w/index.php?title=File:DyFibers2.jpg License: Public Domain Contributors: Clement Yonker; PNNL (US department of Energy)
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9
Holmium
1
Holmium 67
dysprosium ← holmium → erbium
↑
Ho ↓
Es Periodic Table - Extended Periodic Table
General Name, Symbol, Number
holmium, Ho, 67
Element category
lanthanides
Group, Period, Block
n/a, 6, f
Appearance
silvery white
Standard atomic weight
164.93032(2) g·mol
Electron configuration
[Xe] 4f
−1
11
Electrons per shell
2
6s
2, 8, 18, 29, 8, 2 Physical properties
Phase
solid
Density (near r.t.)
8.79 g·cm−3
Liquid density at m.p.
8.34 g·cm−3
Melting point
1734 K (1461 °C, 2662 °F)
Boiling point
2993 K (2720 °C, 4928 °F)
Heat of fusion
17.0 kJ·mol−1
Heat of vaporization
265 kJ·mol−1
Specific heat capacity
(25 °C) 27.15 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1432
1584
(1775)
(2040)
(2410)
(2964)
Atomic properties
Holmium
2
Crystal structure
hexagonal
Oxidation states
3 (basic oxide)
Electronegativity
1.23 (Pauling scale)
Ionization energies (more)
1st: 581.0 kJ·mol−1 2nd: 1140 kJ·mol−1 3rd: 2204 kJ·mol−1
Atomic radius
176 pm
Covalent radius
192±7 pm Miscellaneous
Magnetic ordering
paramagnetic
Electrical resistivity
(r.t.) (poly) 814 nΩ·m
Thermal conductivity
(300 K) 16.2 W·m
Thermal expansion
(r.t.) (poly) 11.2 µm/(m·K)
Speed of sound (thin rod)
(20 °C) 2760 m/s
Young's modulus
64.8 GPa
Shear modulus
26.3 GPa
Bulk modulus
40.2 GPa
Poisson ratio
−1
·K
−1
0.231
Vickers hardness
481 MPa
Brinell hardness
746 MPa
CAS registry number
7440-60-0 Most-stable isotopes
Main article: Isotopes of holmium iso
NA
half-life
DM
DE (MeV)
DP
163
syn
4570 yr
ε
0.003
163
164
syn
29 min
ε
0.987
164
Ho Ho
165
Ho
100%
Dy Dy
165
Ho is stable with 98 neutron
166
syn
26.763 h
β−
1.855
166
167
syn
3.1 h
β−
1.007
167
Ho Ho
Er Er
References
Holmium (pronounced /ˈhoʊlmiəm/) is a chemical element with the symbol Ho and atomic number 67. Part of the lanthanide series, holmium is a relatively soft and malleable silvery-white metallic element, which is stable in dry air at room temperature. A rare earth metal, it is found in the minerals monazite and gadolinite. Holmium has the highest magnetic strength of any element and therefore is used for the polepieces of the strongest
Holmium
3
static magnets. Because holmium can absorb nuclear fission-bred neutrons, it is also used in nuclear control rods.
Characteristics Physical Holmium is a relatively soft and malleable element that is fairly corrosion-resistant and stable in dry air at standard temperature and pressure. In moist air and at higher temperatures, however, it quickly oxidizes, forming a yellowish oxide. In pure form, holmium possesses a metallic, bright silvery luster.
Ho2O3, left: natural light, right: fluorescent lamp light
Holmium oxide has some fairly dramatic color changes depending on the lighting conditions. In daylight, it is a tannish yellow color. Under trichromatic light, it is a fiery orange red, almost indistinguishable from the way erbium oxide looks under this same lighting. This is related to the sharp emission bands of the phosphors.[1]
A trivalent metallic rare earth element, holmium has the highest magnetic moment (10.6 µB) of any naturally-occurring element and possesses other unusual magnetic properties. When combined with yttrium, it forms highly magnetic compounds.[2]
Chemical Holmium metal tarnishes slowly in air and burns readily to form holmium(III) oxide: 4 Ho + 3 O2 → 2 Ho2O3 Holmium is quite electropositive and reacts slowly with cold water and quite quickly with hot water to form holmium hydroxide: 2 Ho (s) + 6 H2O (g) → 2 Ho(OH)3 (aq) + 3 H2 (g) Holmium metal reacts with all the halogens: 2 Ho (s) + 3 F2 (g) → 2 HoF3 (s) [pink] 2 Ho (s) + 3 Cl2 (g) → 2 HoCl3 (s) [yellow] 2 Ho (s) + 3 Br2 (g) → 2 HoBr3 (s) [yellow] 2 Ho (s) + 3 I2 (g) → 2 HoI3 (s) [yellow] Holmium dissolves readily in dilute sulphuric acid to form solutions containing the yellow Ho(III) ions, which exist as a [Ho(OH2)9]3+ complexes:[3] 2 Ho (s) + 3 H2SO4 (aq) → 2 Ho3+ (aq) + 3 SO2−4 (aq) + 3 H2 (g)
Holmium
Isotopes Natural holmium contains one stable isotope, holmium-165. Some synthetic radioactive isotopes are known, the most stable one is holmium-163, with a half life of 4570 years. All other radioisotopes have ground-state half lives not greater than 1.117 days, and most have half lives under 3 hours. However, the metastable 166m1Ho has a half life of around 1200 years because of its high spin. This fact, combined with a high excitation energy resulting in a particularly rich spectrum of decay gamma rays produced when the metastable state de-excites, makes this isotope useful in nuclear physics experiments as a means for calibrating energy responses and intrinsic efficiencies of gamma ray spectrometers.
History Holmium (Holmia, Latin name for Stockholm) was discovered by Marc Delafontaine and Jacques-Louis Soret in 1878 who noticed the aberrant spectrographic absorption bands of the then-unknown element (they called it "Element X").[4] [5] Later in 1878, Per Teodor Cleve independently discovered the element while he was working on erbia earth (erbium oxide).[6] [7] Using the method developed by Carl Gustaf Mosander, Cleve first removed all of the known contaminants from erbia. The result of that effort was two new materials, one brown and one green. He named the brown substance holmia (after the Latin name for Cleve's home town, Stockholm) and the green one thulia. Holmia was later found to be the holmium oxide and thulia was thulium oxide.[8]
Occurrence and production Like all other rare earths, holmium is not naturally found as a free element. It does occur combined with other elements in the minerals gadolinite, monazite, and in other rare-earth minerals. The main mining areas are China, United States, Brazil, India, Sri Lanka and Australia with reserves of holmium estimated as 400,000 tonnes.[8] It is commercially extracted via ion-exchange from monazite sand (0.05% holmium) but is still difficult to Gadolinite separate from other rare earths. The element has been isolated through the reduction of its anhydrous chloride or fluoride with metallic calcium.[9] Its estimated abundance in the Earth's crust is 1.3 mg/kg. Holmium obeys the Oddo-Harkins rule: as an odd-numbered element, it is less abundant than its immediate even numbered neighbors, dysprosium and erbium. However, it is the most abundant of the odd-numbered heavy lanthanides. The principal current source are some of the ion-adsorption clays of southern China. Some of these have a rare-earth composition similar to that found in xenotime or gadolinite. Yttrium makes up about two-thirds of the total by weight; holmium is around 1.5%. The original ores themselves are very lean, maybe only 0.1% total lanthanide, but are easily extracted.[10] Holmium is relatively inexpensive for a rare-earth metal with the price about US$ 1000 per kg.[11]
4
Holmium
Applications Holmium has the highest magnetic strength of any element and therefore is used to create the strongest artificially-generated magnetic fields, when placed within high-strength magnets as a magnetic pole piece (also called a magnetic flux concentrator). Since it can absorb nuclear fission-bred neutrons, it is also used in nuclear control rods.[8] A solution of 4% holmium oxide in 10% Holmium is used in yttrium-iron-garnet (YIG) and perchloric acid, permanently fused into yttrium-lanthanum-fluoride (YLF) solid state lasers a quartz cuvette as an optical found in microwave equipment (which are in turn found calibration standard in a variety of medical and dental settings). Holmium lasers emit at 2.08 microns and therefore are safe to eyes. They are used in medical, dental and fiber-optical applications.[2]
Holmium is one of the colorants used for cubic zirconia and glass, providing yellow or red coloring.[12] Glass containing holmium oxide and holmium oxide solutions (usually in perchloric acid) have sharp optical absorption peaks in the spectral range 200-900 nm. They are therefore used as a calibration standard for optical spectrophotometers[13] and are available commercially.[14] The radioactive but long-lived Ho-166m1 (see "Isotopes" above) is used in calibration of gamma ray spectrometers.[15]
Precautions The element, as with other rare earths, appears to have a low degree of acute toxicity. Holmium plays no biological role in humans but may be able to stimulate metabolism.[9]
See also • Holmium compounds
References [1] Su, Yiguo (2008). "Hydrothermal Synthesis of GdVO4:Ho3+ Nanorods with a Novel White-light Emission". Chemistry Letters 37: 762. doi: 10.1246/cl.2008.762 (http:/ / dx. doi. org/ 10. 1246/ cl. 2008. 762). [2] C. K. Gupta, Nagaiyar Krishnamurthy (2004). Extractive metallurgy of rare earths (http:/ / books. google. co. jp/ books?id=F0Bte_XhzoAC& pg=PA32). CRC Press. p. 32. ISBN 0415333407. . [3] " Chemical reactions of Holmium (https:/ / www. webelements. com/ holmium/ chemistry. html)". Webelements. . Retrieved 2009-06-06. [4] Jacques-Louis Soret (1878). " Sur les spectres d'absorption ultra-violets des terres de la gadolinite (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k3043m/ f1124. table)". Comptes rendus de l'Académie des sciences 87: 1062. . [5] Jacques-Louis Soret (1879). " Sur le spectre des terres faisant partie du groupe de l'yttria (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k3046j/ f550. table)". Comptes rendus de l'Académie des sciences 89: 521. . [6] Per Teodor Cleve (1879). " Sur deux nouveaux éléments dans l'erbine (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k3046j/ f432. table)". Comptes rendus de l'Académie des sciences 89: 478. . [7] Per Teodor Cleve (1879). " Sur l'erbine (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k3046j/ f759. table)". Comptes rendus de l'Académie des sciences 89: 708. . [8] John Emsley (2001). Nature's building blocks: an A-Z guide to the elements (http:/ / books. google. com/ books?id=Yhi5X7OwuGkC& pg=PA181). US: Oxford University Press. pp. 181–182. ISBN 0198503415. . [9] C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304814.
5
Holmium [10] Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds (http:/ / books. google. com/ books?id=Xqj-TTzkvTEC& pg=PA338). McGraw-Hill. pp. 338–339. ISBN 0070494398. . Retrieved 2009-06-06. [11] James B. Hedrick. " Rare-Earth Metals (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ rare_earths/ 740798. pdf)". USGS. . Retrieved 2009-06-06. [12] " Cubic zirconia (http:/ / www. geologyrocks. co. uk/ tutorials/ cubic_zirconia)". . Retrieved 2009-06-06. [13] R. P. MacDonald (1964). " Uses for a Holmium Oxide Filter in Spectrophotometry (http:/ / www. clinchem. org/ cgi/ reprint/ 10/ 12/ 1117. pdf)". Clinical Chemistry 10: 1117. . [14] " Holmium Glass Filter for Spectrophotometer Calibration (http:/ / www. labshoponline. com/ holmium-glass-filter-spectrophotometer-calibration-p-88. html)". . Retrieved 2009-06-06. [15] Ming-Chen Yuan, Jeng-Hung Lee and Wen-Song Hwang (2002). "The absolute counting of 166mHo, 58Co and 88 Y". Applied Radiation and Isotopes 56: 424. doi: 10.1016/S0969-8043(01)00226-3 (http:/ / dx. doi. org/ 10. 1016/ S0969-8043(01)00226-3).
• Guide to the Elements – Revised Edition, Albert Stwertka, (Oxford University Press; 1998) ISBN 0-19-508083-1
External links • WebElements.com – Holmium (http:/ / www. webelements. com/ webelements/ elements/ text/ Ho/ index. html) (also used as a reference) • American Elements – Holmium (http:/ / www. americanelements. com/ hoinfo3. htm) (also used as a reference)
6
Article Sources and Contributors
Article Sources and Contributors Holmium Source: http://en.wikipedia.org/w/index.php?oldid=307225762 Contributors: ABF, AWeishaupt, Abu-Fool Danyal ibn Amir al-Makhiri, Addshore, Ahoerstemeier, AlimanRuna, Archimerged, Arkuat, AzaToth, Bananamanttm69, Beetstra, Bill-on-the-Hill, Bimston, BlueEarth, Borislav Dopudja, Bryan Derksen, CanisRufus, Century0, Conversion script, Cool3, Cygfrydd Llewellyn, Cyp, Darrien, David Latapie, Dbenbenn, Donarreiskoffer, DrBob, Drmies, Dysprosia, Edgar181, Emperorbma, FaerieInGrey, Femto, Geneb1955, Grendelkhan, Hak-kâ-ngìn, Helge Skjeveland, HexaChord, IWhisky, Icairns, Ideyal, IvanLanin, J.delanoy, Jaan513, Jaraalbe, Jdrewitt, Joanjoc, Karl-Henner, Kwamikagami, LA2, LilHelpa, Looxix, Loren.wilton, Marc Venot, Marsa Lahminal, Materialscientist, Mav, MichaelMaggs, Minesweeper, Nergaal, Nihiltres, Pedant, Plexust, Poolkris, Quadell, RTC, Remember, Riana, Roberta F., Romanm, Rursus, Saperaud, Sbharris, Schneelocke, Sengkang, Sl, Smallweed, Steve Hart, Stifynsemons, Stone, THEN WHO WAS PHONE?, Tagishsimon, Tetracube, Tevildo, The Firewall, Versus22, Vsmith, Wrenchelle, Yekrats, Yyy, 134 anonymous edits
Image Sources, Licenses and Contributors image:Ho-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Ho-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Paddy, Saperaud Image: Holmium 1-croprotated.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Holmium_1-croprotated.jpg License: unknown Contributors: User:Materialscientist Image:Holmium(III) oxide.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Holmium(III)_oxide.jpg License: unknown Contributors: User:Filousoph File:Gadolinitas.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Gadolinitas.jpg License: GNU Free Documentation License Contributors: Original uploader was WesternDevil at lt.wikipedia File:HoOxideSolution.jpg Source: http://en.wikipedia.org/w/index.php?title=File:HoOxideSolution.jpg License: Public Domain Contributors: NIST
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
7
Erbium
1
Erbium 68
holmium ← erbium → thulium
↑
Er ↓
Fm Periodic Table - Extended Periodic Table
General Name, Symbol, Number
erbium, Er, 68
Element category
lanthanides
Group, Period, Block
n/a, 6, f
Appearance
silvery white
Standard atomic weight
167.259(3) g·mol
Electron configuration
[Xe] 4f
−1
12
Electrons per shell
2
6s
2, 8, 18, 30, 8, 2 Physical properties
Phase
solid
Density (near r.t.)
9.066 g·cm−3
Liquid density at m.p.
8.86 g·cm−3
Melting point
1802 K (1529 °C, 2784 °F)
Boiling point
3141 K (2868 °C, 5194 °F)
Heat of fusion
19.90 kJ·mol−1
Heat of vaporization
280 kJ·mol−1
Specific heat capacity
(25 °C) 28.12 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1504
1663
(1885)
(2163)
(2552)
(3132)
Erbium
2 Atomic properties
Crystal structure Oxidation states Electronegativity Ionization energies (more)
hexagonal 3 (basic oxide) 1.24 (Pauling scale) 1st: 589.3 kJ·mol−1 2nd: 1150 kJ·mol−1 3rd: 2194 kJ·mol−1
Atomic radius
176 pm
Covalent radius
189±6 pm Miscellaneous
Magnetic ordering
paramagnetic at 300 K
Electrical resistivity
(r.t.) (poly) 0.860 µΩ·m
Thermal conductivity
(300 K) 14.5 W·m
Thermal expansion
(r.t.) (poly) 12.2 µm/(m·K)
Speed of sound (thin rod)
(20 °C) 2830 m/s
Young's modulus
69.9 GPa
Shear modulus
28.3 GPa
Bulk modulus
44.4 GPa
Poisson ratio
0.237
Vickers hardness
589 MPa
Brinell hardness
814 MPa
CAS registry number
7440-52-0 Most-stable isotopes
−1
·K
−1
Erbium
3
Main article: Isotopes of erbium iso 160
Er
NA syn
half-life 28.58 h
162
0.139%
162
164
1.601%
164
Er Er
165
Er
syn
10.36 h 166
167
22.869%
167
168
26.978%
168
Er
169
Er
170
Er
syn 14.910%
0.330
160
0.376
165
0.351
169
−
1.490
171
−
0.891
172
ε
Ho
Er is stable with 99 neutron Er is stable with 100 neutron
9.4 d
−
β
Tm
170
Er is stable with 102 neutron
syn
7.516 h
β
172
syn
49.3 h
β
Er
Ho
Er is stable with 98 neutron
171
Er
DP
Er is stable with 96 neutron
33.503%
Er
ε
DE (MeV)
Er is stable with 94 neutron
166
Er
DM
Tm Tm
References
Erbium (pronounced /ˈɜrbiəm/) is a chemical element with the symbol Er and atomic number 68. A rare, silvery, white metallic lanthanide, erbium is solid in its normal state. It is a rare earth element associated with several other rare elements in the mineral gadolinite from Ytterby in Sweden.
Characteristics Physical A trivalent element, pure erbium metal is malleable (or easily shaped), soft yet stable in air, and does not oxidize as quickly as some other rare-earth metals. Its salts are rose-colored, and the element has characteristic sharp absorption spectra bands in visible light, ultraviolet, and near infrared. Otherwise it looks much like the other rare earths. Its sesquioxide is called erbia. Erbium's properties are to a degree dictated by the kind and amount of impurities present. Erbium does not play any known biological role, but is thought to be able to stimulate metabolism.[1] Erbium-doped glasses or crystals can be used as optical amplification media, where erbium ions are optically pumped at around 980 nm or 1480 nm and then radiate light at 1550 nm. This process can be used to create lasers and optical amplifiers. The 1550 nm wavelength is especially important for optical communications because standard single mode optical fibers have minimal loss at this particular wavelength. A large variety of medical applications can be found (i.e. dermatology, dentistry) by utilizing the 2940 nm emission (see Er:YAG laser) which is highly absorbed in water (absorption coefficient about 12000/cm). Erbium is ferromagnetic below 19 K, antiferromagnetic between 19 and 80 K and paramagnetic above 80 K [2] .
Erbium
4
Erbium can form propeller-shaped atomic clusters Er3N, where the distance between the erbium atoms is 0.35 nm. Those clusters can be isolated by encapsulating them into fullerene molecules, as confirmed by transmission electron microscopy.[3]
Chemical Erbium metal tarnishes slowly in air and burns readily to form erbium(III) oxide: 4 Er + 3 O2 → 2 Er2O3 Erbium is quite electropositive and reacts slowly with cold water and quite quickly with hot water to form erbium hydroxide: 2 Er (s) + 6 H2O (g) → 2 Er(OH)3 (aq) + 3 H2 (g) Erbium metal reacts with all the halogens: 2 Er (s) + 3 F2 (g) → 2 ErF3 (s) [pink] 2 Er (s) + 3 Cl2 (g) → 2 ErCl3 (s) [violet] 2 Er (s) + 3 Br2 (g) → 2 ErBr3 (s) [violet] 2 Er (s) + 3 I2 (g) → 2 ErI3 (s) [violet] Erbium dissolves readily in dilute sulphuric acid to form solutions containing the yellow Er(III) ions, which exist as a [Er(OH2)9]3+ complexes:[4] 2 Er (s) + 3 H2SO4 (aq) → 2 Er3+ (aq) + 3 SO2−4 (aq) + 3 H2 (g)
Isotopes Naturally occurring erbium is composed of 6 stable isotopes, Er-162, Er-164, Er-166, Er-167, Er-168, and Er-170 with Er-166 being the most abundant (33.503% natural abundance). 29 radioisotopes have been characterized, with the most stable being Er-169 with a half-life of 9.4 days, Er-172 with a half-life of 49.3 hours, Er-160 with a half-life of 28.58 hours, Er-165 with a half-life of 10.36 hours, and Er-171 with a half-life of 7.516 hours. All of the remaining radioactive isotopes have half-lives that are less than 3.5 hours, and the majority of these have half-lives that are less than 4 minutes. This element also has 13 meta states, with the most stable being Er-167m (t½ 2.269 seconds).[5] The isotopes of erbium range in atomic weight from 142.9663 u (Er-143) to 176.9541 u (Er-177). The primary decay mode before the most abundant stable isotope, Er-166, is electron capture, and the primary mode after is beta decay. The primary decay products before Er-166 are element 67 (holmium) isotopes, and the primary products after are element 69 (thulium) isotopes.[5]
History Erbium (for Ytterby, a town in Sweden) was discovered by Carl Gustaf Mosander in 1843. Mosander separated "yttria" from the mineral gadolinite into three fractions which he called yttria, erbia, and terbia. He named the new element after the town of Ytterby where large concentrations of yttria and erbium are located. Erbia and terbia, however, were confused at this time. After 1860, terbia was renamed erbia and after 1877 what had been known as erbia was renamed terbia. Fairly pure Er2O3 was independently isolated in 1905 by Georges Urbain and Charles James. Reasonably pure metal wasn't produced until 1934 when Klemm and Bommer reduced the anhydrous chloride with potassium vapor. It was only in the 1990s that the price for Chinese-derived erbium oxide became low enough for
Erbium erbium to be considered for use as a colorant in art glass.[6]
Occurrence Like other rare earths, this element is never found as a free element in nature but is found bound in monazite sand ores. It has historically been very difficult and expensive to separate rare earths from each other in their ores but ion-exchange production techniques developed in the late 20th century have greatly brought down the cost of production of all rare-earth metals and their chemical compounds. The principal commercial sources of erbium are from the minerals xenotime and Monazite sand euxenite, and most recently, the ion adsorption clays of southern China. In the high-yttrium versions of these ore concentrates, yttrium is about two-thirds of the total by weight, and erbia is about 4-5%. This is enough erbium to impart a distinct pink color to the solution when the concentrate is dissolved in acid. This color behavior is highly similar to what Mosander and the other early workers in the lanthanides would have seen, in their extracts from Ytterby gadolinite. The concentration of erbium in the Earth crust is about 2.8 mg/kg and in the sea water 0.9 ng/L.[7]
Production Crushed minerals are attacked by hydrochloric or sulfuric acid that transforms insoluble rare-earth oxides into soluble chlorides or sulfates. The acidic filtrates are partially neutralized with caustic soda to pH 3-4. Thorium precipitates out of solution as hydroxide and is removed. After that the solution is treated with ammonium oxalate to convert rare earths in to their insoluble oxalates. The oxalates are converted to oxides by annealing. The oxides are dissolved in nitric acid that excludes one of the main components, cerium, whose oxide is insoluble in HNO3. The solution is treated with magnesium nitrate to produce a crystallized mixture of double salts of rare-earth metals. The salts are separated by ion exchange. In this process, rare-earth ions are sorbed onto suitable ion-exchange resin by exchange with hydrogen, ammonium or cupric ions present in the resin. The rare earth ions are then selectively washed out by suitable complexing agent.[7] Erbium metal is obtained from its oxide or salts by heating with calcium at 1450 °C under argon atmosphere.[7]
5
Erbium
Applications Erbium's everyday uses are varied. It is commonly used as a photographic filter, and because of its resilience it is useful as a metallurgical additive. Other uses: • Used in nuclear technology as a nuclear poison, as in neutron-absorbing control rods.[1] [8]
• When added to vanadium as an alloy, erbium lowers hardness and improves workability.[9] • Erbium oxide has a pink color, and is sometimes used as a colorant for glass, cubic zirconia and porcelain. The glass is then often used in sunglasses and cheap jewelry.[9] • Erbium-doped optical silica-glass fibers are the active element in erbium-doped fiber amplifiers (EDFAs), which are widely used in optical communications.[10] The same fibers can be used to create fiber lasers. Co-doping of optical fiber with Er and Yb is used in high-power Er/Yb fiber lasers, which gradually replace CO2 lasers for metal welding and cutting applications. Erbium can also be used in erbium-doped waveguide amplifiers.[1] • An erbium-nickel alloy Er3Ni has an unusually high specific heat capacity at liquid-helium temperatures and is used in cryocoolers; a mixture of 65% Er3Co and 35% Er0.9Yb0.1Ni by volume improves the specific heat capacity even more.[11] [12]
Precautions As with the other lanthanides, erbium compounds are of low to moderate toxicity, although their toxicity has not been investigated in detail. Metallic erbium in dust form presents a fire and explosion hazard.
See also • • • •
Erbium compounds Terbium Ytterbium Yttrium
Further reading • Guide to the Elements – Revised Edition, Albert Stwertka, (Oxford University Press; 1998) ISBN 0-19-508083-1
External links • WebElements.com – Erbium [13] (also used as a reference) • It's Elemental – Erbium [14] • Chemical Elements: Erbium http:/ / www. chemicalelements. com/ elements/ er. html pnb:میبرا
6
Erbium
References [1] Emsley, John (2001). " Erbium (http:/ / books. google. com/ books?id=j-Xu07p3cKwC)". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. pp. 136–139. ISBN 0-19-850340-7. . [2] M. Jackson (2000). " Magnetism of Rare Earth (http:/ / www. irm. umn. edu/ quarterly/ irmq10-3. pdf)". The IRM quarterly 10 (3): 1. . [3] Yuta Sato,; Kazu Suenaga,; Shingo Okubo; Toshiya Okazaki,; Sumio Iijima (2007). "Structures of D5d-C80 and Ih-Er3N@C80 Fullerenes and Their Rotation Inside Carbon Nanotubes Demonstrated by Aberration-Corrected Electron Microscopy". Nano Letters 7: 3704. doi: 10.1021/nl0720152 (http:/ / dx. doi. org/ 10. 1021/ nl0720152). [4] " Chemical reactions of Erbium (https:/ / www. webelements. com/ erbium/ chemistry. html)". Webelements. . Retrieved 2009-06-06. [5] Georges, Audi (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3–128. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [6] Aaron John Ihde (1984). The development of modern chemistry (http:/ / books. google. com/ books?id=34KwmkU4LG0C& pg=PA377& ). Courier Dover Publications. p. 378-379. ISBN 0486642356. . [7] Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds (http:/ / books. google. com/ books?id=Xqj-TTzkvTEC& pg=PA293). McGraw-Hill. pp. 293–295. ISBN 0070494398. . Retrieved 2009-06-06. [8] edited by Theodore A. Parish, Vyacheslav V. Khromov, Igor Carron. (1999). " Use of UraniumErbium and PlutoniumErbium Fuel in RBMK Reactors (http:/ / books. google. de/ books?id=aamn7uifb3gC)". Safety issues associated with Plutonium involvement in the nuclear fuel cycle. CBoston: Kluwer. pp. 121–125. ISBN 9780792355939. . [9] C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304814. [10] P.C. Becker, N.A. Olsson, J.R. Simpson ; (1999). Erbium-doped fiber amplifiers fundamentals and technology (http:/ / books. google. com/ books?hl=de& lr=& id=uAOq75yt5CcC). San Diego: Academic Press. ISBN 9780120845903. . [11] Peter Kittel, ed. Advances in Cryogenic Engineering volume 39a. [12] Ackermann, Robert A. (1997). Cryogenic Regenerative Heat Exchangers (http:/ / books. google. com/ books?id=nIzviZ_-_NsC). Springer. p. 58. ISBN 9780306454493. . [13] http:/ / www. webelements. com/ webelements/ elements/ text/ Er/ index. html [14] http:/ / education. jlab. org/ itselemental/ ele068. html
7
Article Sources and Contributors
Article Sources and Contributors Erbium Source: http://en.wikipedia.org/w/index.php?oldid=308121748 Contributors: 2over0, A new name 2008, AWeishaupt, Agateller, Ahoerstemeier, AlimanRuna, Archimerged, Arkuat, Beetstra, Bergsten, Betacommand, Bettia, BillC, BlueEarth, Bobo192, Bryan Derksen, Californium-256, Cmdrjameson, Conversion script, Coppertwig, DMacks, Darrien, David Latapie, Donarreiskoffer, DrBob, Easwarno1, Eddideigel, Edgar181, El C, Elf, Emperorbma, Encyclopedia77, Femto, Fivemack, Flava flav snap it, Flowanda, Geni, Goatasaur, Grendelkhan, Hak-kâ-ngìn, Hankwang, Heron, Icairns, Ideyal, IvanLanin, J.delanoy, Jaan513, Jaraalbe, Jerry teps, Jessesaurus, Karelj, Karl-Henner, Kristaga, Kwamikagami, LA2, Leafyplant, Lexicon, Lightmouse, Lukay 79, Magnus Manske, Marc Venot, Materialscientist, Mav, Mentifisto, Minesweeper, MxM, Nergaal, Nick Y., Nihiltres, Pharaoh of the Wizards, Plexust, Poolkris, Quadell, QueenCake, RTC, Remember, Res2216firestar, Reza kalani, Roberta F., Romanm, Saga City, Saperaud, Schneelocke, Sl, Srleffler, Steve Hart, Stifynsemons, Stone, Suisui, Tagishsimon, Tetracube, The Anome, Turk oğlan, UnHoly, Versus22, Vicki Rosenzweig, Vsmith, Warut, Wikip rhyre, Wiwaxia, Wtmitchell, Yekrats, Yyy, 113 anonymous edits
Image Sources, Licenses and Contributors image:Er-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Er-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Paddy, Saperaud Image: Erbium-crop.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Erbium-crop.jpg License: unknown Contributors: User:Materialscientist File:MonaziteUSGOV.jpg Source: http://en.wikipedia.org/w/index.php?title=File:MonaziteUSGOV.jpg License: Public Domain Contributors: Saperaud
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
8
Thulium
1
Thulium 69
erbium ← thulium → ytterbium
↑
Tm ↓
Md Periodic Table - Extended Periodic Table
General Name, Symbol, Number
thulium, Tm, 69
Element category
lanthanides
Group, Period, Block
n/a, 6, f
Appearance
silvery gray
Standard atomic weight
168.93421(2) g·mol
Electron configuration
[Xe] 4f
−1
13
Electrons per shell
2
6s
2, 8, 18, 31, 8, 2 Physical properties
Phase
solid
Density (near r.t.)
9.32 g·cm−3
Liquid density at m.p.
8.56 g·cm−3
Melting point
1818 K (1545 °C, 2813 °F)
Boiling point
2223 K (1950 °C, 3542 °F)
Heat of fusion
16.84 kJ·mol−1
Heat of vaporization
247 kJ·mol−1
Specific heat capacity
(25 °C) 27.03 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1117
1235
1381
1570
(1821)
(2217)
Atomic properties Crystal structure
hexagonal
Thulium
2
Oxidation states
2, 3, 4 (basic oxide)
Electronegativity
1.25 (Pauling scale)
Ionization energies (more)
1st: 596.7 kJ·mol−1 2nd: 1160 kJ·mol−1 3rd: 2285 kJ·mol−1
Atomic radius
176 pm
Covalent radius
190±10 pm Miscellaneous
Magnetic ordering
paramagnetic at 300 K
Electrical resistivity
(r.t.) (poly) 676 nΩ·m
Thermal conductivity
(300 K) 16.9 W·m
Thermal expansion
(r.t.) (poly) 13.3 µm/(m·K)
Young's modulus
74.0 GPa
Shear modulus
30.5 GPa
Bulk modulus
44.5 GPa
−1
Poisson ratio
−1
·K
0.213
Vickers hardness
520 MPa
Brinell hardness
471 MPa
CAS registry number
7440-30-4 Most-stable isotopes
Main article: Isotopes of thulium iso
NA
half-life
DM
DE (MeV)
DP
167
syn
9.25 d
ε
0.748
167
168
syn
93.1 d
ε
1.679
168
Tm Tm
169
Tm
100%
Er Er
169
Tm is stable with 100 neutron
170
syn
128.6 d
β−
0.968
170
171
syn
1.92 y
β−
0.096
171
Tm Tm
Yb Yb
References
Thulium (pronounced /ˈθjuːliəm/) is a chemical element that has the symbol Tm and atomic number 69. Thulium is the least abundant of the lanthanides (promethium is less abundant than thulium, but it is not found naturally on Earth). It is an easily workable metal with a bright silvery-gray luster and can be cut by a knife. Despite its high price and rarity, thulium is used as radiation source in portable X-ray devices and in solid-state lasers.
Thulium
3
Properties Physical Pure thulium metal has a bright, silvery luster. It is reasonably stable in air, but should be protected from moisture. The metal is soft, malleable, and ductile, and can be cut with a knife.[1] Thulium is ferromagnetic below 32 K, antiferromagnetic between 32 and 56 K and paramagnetic above 56 K [2] .
Chemical Thulium metal tarnishes slowly in air and burns readily at 150 °C to form thulium(III) oxide: 4 Tm + 3 O2 → 2 Tm2O3 Thulium is quite electropositive and reacts slowly with cold water and quite quickly with hot water to form thulium hydroxide: 2 Tm (s) + 6 H2O (g) → 2 Tm(OH)3 (aq) + 3 H2 (g) Thulium reacts with all the halogens at temperatures. Reactions are slow at room temperature, but are vigorous above 200 °C: 2 Tm (s) + 3 F2 (g) → 2 TmF3 (s) [white] 2 Tm (s) + 3 Cl2 (g) → 2 TmCl3 (s) [yellow] 2 Tm (s) + 3 Br2 (g) → 2 TmBr3 (s) [white] 2 Tm (s) + 3 I2 (g) → 2 TmI3 (s) [yellow] Thulium dissolves readily in dilute sulfuric acid to form solutions containing the pale green Tm(III) ions, which exist as a [Tm(OH2)9]3+ complexes:[3] 2 Tm (s) + 3 H2SO4 (aq) → 2 Tm3+ (aq) + 3 SO2−4 (aq) + 3 H2 (g)
Thulium reacts with various metallic and non-metallic elements forming a range of binary compounds, including TmN, TmS, TmC2, Tm2C3, TmH2, TmH3, TmSi2, TmGe3, TmB4, TmB6 and TmB12. In those compounds, thulium exhibits valence states +2, +3 and +4, however, the +3 state is most common and only this state has been observed in Tm solutions.[4]
Isotopes Naturally occurring thulium is composed of one stable isotope, Tm-169 (100% natural abundance). Thirty one radioisotopes have been characterized, with the most stable being Tm-171 with a half-life of 1.92 years, Tm-170 with a half-life of 128.6 days, Tm-168 with a half-life of 93.1 days, and Tm-167 with a half-life of 9.25 days. All of the remaining radioactive isotopes have half-lives that are less than 64 hours, and the majority of these have half-lives that are less than 2 minutes. This element also has 14 meta states, with the most stable being Tm-164m (t½ 5.1 minutes), Tm-160m (t½ 74.5 seconds) and Tm-155m (t½ 45 seconds). The isotopes of thulium range in atomic weight from 145.966 u (Tm-146) to 176.949 u (Tm-177). The primary decay mode before the most abundant stable isotope, Tm-169, is electron capture, and the primary mode after is beta emission. The primary decay products before Tm-169 are element 68 (erbium) isotopes, and the primary products after are element 70 (ytterbium) isotopes.[5]
Thulium
History Thulium was discovered by Swedish chemist Per Teodor Cleve in 1879 by looking for impurities in the oxides of other rare earth elements (this was the same method Carl Gustaf Mosander earlier used to discover some other rare earth elements). Cleve started by removing all of the known contaminants of erbia (Er2O3). Upon additional processing, he obtained two new substances; one brown and one green. The brown substance turned out to be the oxide of the element holmium and was named holmia by Cleve, and the green substance was the oxide of an unknown element. Cleve named the oxide thulia and its element thulium after Thule, Scandinavia. In 1911, Theodore William Richards performed 15,000 recrystallizations of thulium bromate to obtain pure sample of the element and so to accurately determine its atomic weight.[6] Thulium was so rare that none of the early workers had enough of it to purify sufficiently to actually see the green color; they had to be content with spectroscopically observing the strengthening of the two characteristic absorption bands, as erbium was progressively removed. The first researcher to obtain nearly pure thulium was Charles James, a British expatriate working on a large scale at New Hampshire College in Durham NH. In 1911 he reported his results, having used his discovered method of bromate fractional crystallization to do the purification. He famously needed 15,000 "operations" to establish that the material was homogeneous.[7] High-purity thulium oxide was first offered commercially in the late 1950s, as a result of the adoption of ion-exchange separation technology. Lindsay Chemical Division of American Potash & Chemical Corporation offered it in grades of 99% and 99.9% purity. The price has oscillated between US$ 4,600 and 13,300 in the period from 1959 to 1998 for 99.9% purity, and it was second highest for lanthanides behind lutetium.[8] [9]
Occurrence and production The element is never found in nature in pure form, but it is found in small quantities in minerals with other rare earths. Its abundance in the Earth crust is 0.5 mg/kg.[6] Thulium is principally extracted from monazite (~0.007% thulium) ores found in river sands, through ion-exchange. Newer ion-exchange and solvent-extraction techniques have led to easier separation of the rare earths, which has yielded much lower costs for thulium production. The principal sources today are the ion adsorption clays of southern China. In these, where about two-thirds of the total rare-earth content is yttrium, thulium is about 0.5% (or about tied with lutetium for rarity). The metal can be isolated through reduction of its oxide with lanthanum metal or by calcium reduction in a closed container. None of thulium's natural compounds are commercially important.[1]
4
Thulium
5
Applications Rare and expensive, thulium has few applications:
Laser Holmium-chromium-thulium triple-doped YAG (Ho:Cr:Tm:YAG, or Ho,Cr,Tm:YAG) is an active laser medium material with high efficiency. It lases at 2097 nm and is widely used in military, medicine, and meteorology. Single-element thulium-doped YAG (Tm:YAG) lasers operate between 1930 and 2040 nm.[10] The wavelength of thulium-based lasers is very efficient for superficial ablation of tissue, with minimal coagulation depth in air or in water. This makes thulium lasers attractive for laser-based surgery.[11]
X-ray source Despite its high cost, portable X-ray devices use thulium that has been bombarded in a nuclear reactor as a radiation source. These sources are available for about one year, as tools in medical and dental diagnosis, as well as to detect defects in inaccessible mechanical and electronic components. Such sources do not need extensive radiation protection - only a small cup of lead. [12]
Others Thulium has been used in high temperature superconductors similarly to yttrium. Thulium potentially has use in ferrites, ceramic magnetic materials that are used in microwave equipment.[12]
Biological role and precautions Thulium has no known biological role, although it has been noted that it stimulates metabolism. Soluble thulium salts are regarded as slightly toxic if taken in large amounts, but the insoluble salts are non-toxic. Thulium is not taken up by plant roots to any extend and thus does not get into the human food chain. Vegetables typically contain only one milligram of thulium per tonne (dry weight).[6]
See also • Thulium compounds
External links • WebElements.com – Thulium • It's Elemental – Thulium [14]
[13]
(also used as a reference)
pnb:میلھت
References [1] C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304814. [2] M. Jackson (2000). " Magnetism of Rare Earth (http:/ / www. irm. umn. edu/ quarterly/ irmq10-3. pdf)". The IRM quarterly 10 (3): 1. . [3] " Chemical reactions of Thulium (https:/ / www. webelements. com/ thulium/ chemistry. html)". Webelements. . Retrieved 2009-06-06.
Thulium [4] Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds (http:/ / books. google. com/ books?id=Xqj-TTzkvTEC& pg=PA934). McGraw-Hill. p. 934. ISBN 0070494398. . Retrieved 2009-06-06. [5] Lide, David R. (1998). "Section 11, Table of the Isotopes". Handbook of Chemistry and Physics (87 ed.). Boca Raton, FL: CRC Press. ISBN 0849305942. [6] John Emsley (2001). Nature's building blocks: an A-Z guide to the elements (http:/ / books. google. com/ books?id=Yhi5X7OwuGkC& pg=PA442). US: Oxford University Press. pp. 442-443. ISBN 0198503415. . [7] James, Charles (1911). "Thulium I". J. Am. Chem. Soc. 33 (8): 1332–1344. doi: 10.1021/ja02221a007 (http:/ / dx. doi. org/ 10. 1021/ ja02221a007). [8] James B. Hedrick. " Rare-Earth Metals (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ rare_earths/ 740798. pdf)". USGS. . Retrieved 2009-06-06. [9] Stephen B. Castor and James B. Hedrick. " Rare Earth Elements (http:/ / www. rareelementresources. com/ i/ pdf/ RareEarths-CastorHedrickIMAR7. pdf)". . Retrieved 2009-06-06. [10] Walter Koechner (2006). Solid-state laser engineering (http:/ / books. google. com/ books?id=RK3jK0XWjdMC& pg=PA49& ). Springer. p. 49. ISBN 038729094X. . [11] Frank J. Duarte (2008). Tunable laser applications (http:/ / books. google. com/ books?id=FCDPZ7e0PEgC& pg=PA214& ). CRC Press. p. 214. ISBN 1420060090. . [12] C. K. Gupta, Nagaiyar Krishnamurthy (2004). Extractive metallurgy of rare earths (http:/ / books. google. co. jp/ books?id=F0Bte_XhzoAC& pg=PA32). CRC Press. p. 32. ISBN 0415333407. . [13] http:/ / www. webelements. com/ webelements/ elements/ text/ Tm/ index. html [14] http:/ / education. jlab. org/ itselemental/ ele069. html
6
Article Sources and Contributors
Article Sources and Contributors Thulium Source: http://en.wikipedia.org/w/index.php?oldid=308396809 Contributors: A2Kafir, AWeishaupt, Ahoerstemeier, Alansohn, AlimanRuna, Archimerged, Arex, Baccyak4H, Bryan Derksen, CDN99, Conversion script, Cryptic C62, Darrien, David Latapie, DerHexer, Donarreiskoffer, Dysprosia, Edgar181, El C, Emperorbma, Femto, Flauto Dolce, Greatpatton, Grendelkhan, Gurch, Guyseni, Hak-kâ-ngìn, Heron, Icairns, Icelight, Ideyal, Ilikeverin, IvanLanin, J.delanoy, Jaan513, Jaraalbe, Jcw69, Jediknil, Jimfbleak, Joanjoc, Karelj, Kelovy, Kurykh, Kwamikagami, LA2, LarryMorseDCOhio, Levil, LilHelpa, Lradrama, Marc Venot, Materialscientist, Mav, Michaelbarreto, Midgley, Minesweeper, Nergaal, Netkinetic, Nihiltres, Oosh, Oxymoron83, PeepP, Picapica, Plexust, Poolkris, Quiddity, RTC, Rallette, Remember, RexNL, Reyk, Riana, Rjwilmsi, Roberta F., Romanm, Saperaud, Schneelocke, Shen, Silverhill, Skyramp, Sl, Steve Hart, Stifynsemons, Stone, Svante, Tagishsimon, Tanada, Tetracube, Tonybaloney867, Ugochukwu14, Vsmith, Warut, Yekrats, Yuval madar, Yyy, 133 anonymous edits
Image Sources, Licenses and Contributors image:Tm-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Tm-TableImage.png License: GNU Free Documentation License Contributors: Joolz, Paddy, Paginazero, Saperaud Image: Tm,69.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Tm,69.jpg License: GNU Free Documentation License Contributors: User:RTC Image:Monazit - Madagaskar.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Monazit_-_Madagaskar.jpg License: unknown Contributors: User:Ra'ike
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
7
Ytterbium
1
Ytterbium 70
thulium ← ytterbium → lutetium
↑
Yb ↓
No Periodic Table - Extended Periodic Table
General Name, Symbol, Number
ytterbium, Yb, 70
Element category
lanthanides
Group, Period, Block
n/a, 6, f
Appearance
silvery white
Standard atomic weight
173.04(3) g·mol
Electron configuration
[Xe] 4f
−1
14
Electrons per shell
2
6s
2, 8, 18, 32, 8, 2 Physical properties
Phase
solid
Density (near r.t.)
6.90 g·cm−3
Liquid density at m.p.
6.21 g·cm−3
Melting point
1097 K (824 °C, 1515 °F)
Boiling point
1469 K (1196 °C, 2185 °F)
Heat of fusion
7.66 kJ·mol−1
Heat of vaporization
159 kJ·mol−1
Specific heat capacity
(25 °C) 26.74 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
736
813
910
1047
(1266)
(1465)
Atomic properties Crystal structure
cubic face centered
Ytterbium
Oxidation states Electronegativity Ionization energies (more)
2 3, 2 (basic oxide) ? 1.1 (Pauling scale) 1st: 603.4 kJ·mol−1 2nd: 1174.8 kJ·mol−1 3rd: 2417 kJ·mol−1
Atomic radius
176 pm
Covalent radius
187±8 pm Miscellaneous [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(r.t.) (β, poly) 0.250 µΩ·m
Thermal conductivity
(300 K) 38.5 W·m
Thermal expansion
(r.t.) (β, poly) 26.3 µm/(m·K)
Speed of sound (thin rod)
(20 °C) 1590 m/s
Young's modulus
(β form) 23.9 GPa
Shear modulus
(β form) 9.9 GPa
Bulk modulus
(β form) 30.5 GPa
Poisson ratio
(β form) 0.207
Vickers hardness
206 MPa
Brinell hardness
343 MPa
CAS registry number
7440-64-4 Most-stable isotopes
−1
·K
−1
Ytterbium
3
Main article: Isotopes of ytterbium iso 166
Yb
168
Yb
169
Yb
NA syn 0.13% syn
half-life 56.7 h
32.026 d 170
171
14.28%
171
172
21.83%
172
173
16.13%
173
174
31.83%
174
Yb Yb Yb
175
Yb
176
Yb
177
Yb
syn 12.76% syn
DP
0.304
166
0.909
169
0.470
175
1.399
177
Tm
Yb is stable with 98 neutron
3.04%
Yb
ε
DE (MeV)
168
170
Yb
DM
ε
Tm
Yb is stable with 100 neutron Yb is stable with 101 neutron Yb is stable with 102 neutron Yb is stable with 103 neutron Yb is stable with 104 neutron
4.185 d
−
β
Lu
176
Yb is stable with 106 neutron
1.911 h
−
β
Lu
References
Ytterbium (pronounced /ɨˈtɝːbiəm/) is a chemical element with the symbol Yb and atomic number 70. A soft silvery metallic element, ytterbium is a rare earth of the lanthanide series and is found in the minerals gadolinite, monazite, and xenotime. The element is sometimes associated with yttrium or other related elements and is used in certain steels. Natural ytterbium is a mix of seven stable isotopes. Ytterbium-169, an artificially produced isotope, is used as a gamma ray source.
Characteristics Physical Ytterbium is a soft, malleable and rather ductile element that exhibits a bright silvery luster. A rare earth element, it is easily attacked and dissolved by mineral acids, slowly reacts with water, and oxidizes in air.[2] Ytterbium has three allotropes which are called alpha, beta and gamma and whose transformation points are at −13 °C and 795 °C. The beta form exists at room temperature and has a face-centered crystal structure while the high-temperature gamma form has a body-centered crystal structure.[2] Normally, the beta form has a metallic-like electrical conductivity, but becomes a semiconductor when exposed to around 16,000 atm (1.6 GPa). Its electrical resistivity is tenfold larger at about 39,000 atm (3.9 GPa) but then drops dramatically, to around 10% of its room temperature resistivity value, at 40,000 atm (4 GPa).[2] [3] Contrary to other rare-earth metals, which show antiferromagnetic or/and ferromagnetic ordering at low temperatures, Yb is paramagnetic at any temperatures above 1 K.[4]
Ytterbium
Chemical Ytterbium metal tarnishes slowly in air and burns readily at 200 °C to form ytterbium(III) oxide (Yb2O3) or less stable ytterbium monoxide (YbO). Ytterbium is quite electropositive and reacts slowly with cold water and quite quickly with hot water to form ytterbium hydroxide: 2 Yb (s) + 6 H2O (g) → 2 Yb(OH)3 (aq) + 3 H2 (g) Ytterbium metal reacts with all the halogens: 2 Yb (s) + 3 F2 (g) → 2 YbF3 (s) [white] 2 Yb (s) + 3 Cl2 (g) → 2 YbCl3 (s) [white] 2 Yb (s) + 3 Br2 (g) → 2 YbBr3 (s) [white] 2 Yb (s) + 3 I2 (g) → 2 YbI3 (s) [white] Ytterbium(III) ion absorbs light in the near infrared spectral range, but not in the visible region, so that ytterbia is white, and ytterbium salts of colorless anions are also colorless. Ytterbium dissolves readily in dilute sulfuric acid to form solutions containing the colorless Yb(III) ions, which exist as a [Yb(OH2)9]3+ complexes:[5] 2 Yb (s) + 3 H2SO4 (aq) → 2 Yb3+ (aq) + 3 SO2−4 (aq) + 3 H2 (g)
Compounds Ytterbium shows similar chemical behavior to the rest of the lanthanoid group. Most of the compounds are found in the oxidation state +3, the salts in that oxidation state are nearly colorless. Like europium, samarium or thulium trihalogenes can be reduced by hydrogen or by addition of the metal reduced to the dihalogens, in this case the for example YbCl2. The oxidation state +2 reacts in some behalves similar to the alkaline earth metal compounds, for example the Ytterbium(II) oxide (YbO) shows the same structure than calcium oxide (CaO).[6] • Halides: YbCl2, YbBr3, YbCl3, YbF3 • Oxides: Yb2O3
Isotopes Naturally occurring ytterbium is composed of 7 stable isotopes, Yb-168, Yb-170, Yb-171, Yb-172, Yb-173, Yb-174, and Yb-176, with Yb-174 being the most abundant (31.83% natural abundance). 27 radioisotopes have been characterized, with the most stable being Yb-169 with a half-life of 32.026 days, Yb-175 with a half-life of 4.185 days, and Yb-166 with a half-life of 56.7 hours. All of the remaining radioactive isotopes have half-lives that are less than 2 hours, and the majority of these have half-lives that are less than 20 minutes. This element also has 12 meta states, with the most stable being Yb-169m (t½ 46 seconds). The isotopes of ytterbium range in atomic weight from 147.9674 u (Yb-148) to 180.9562 u (Yb-181). The primary decay mode before the most abundant stable isotope, Yb-174 is electron capture, and the primary mode after is beta emission. The primary decay products before Yb-174 are element 69 (thulium) isotopes, and the primary products after are element 71 (lutetium) isotopes. Of interest to modern quantum optics, the different ytterbium isotopes follow either Bose-Einstein statistics or Fermi-Dirac statistics, leading to interesting behavior in optical lattices.
4
Ytterbium
History Ytterbium was discovered by the Swiss chemist Jean Charles Galissard de Marignac in the year 1878. Marignac found a new component in the earth then known as erbia and named it ytterbia (after Ytterby, the Swedish town where he found the new erbia component). He suspected that ytterbia was a compound of a new element he called ytterbium.[3] In 1907, the French chemist Georges Urbain separated Marignac's ytterbia into two components, neoytterbia and lutecia. Neoytterbia would later become known as the element ytterbium and lutecia would later be known as the element lutetium. Auer von Welsbach independently isolated these elements from ytterbia at about the same time but called them aldebaranium and cassiopeium.[3] The chemical and physical properties of ytterbium could not be determined until 1953 when the first nearly pure ytterbium was produced.[3] The price of ytterbium was relatively stable between 1953 and 1998 at about US$ 1,000/kg.[7]
Occurrence Ytterbium is found with other rare earth elements in several rare minerals. It is most often recovered commercially from monazite sand (0.03% ytterbium). The element is also found in euxenite and xenotime. The main mining areas are China, United States, Brazil, India, Sri Lanka and Australia; and reserves of ytterbium are estimated as about one million tonnes. Ytterbium is normally difficult to separate from other rare earths, but ion-exchange and solvent extraction Euxenite techniques developed in the mid to late 20th century have simplified separation. Known compounds of ytterbium are rare—they haven't been well characterized yet. The abundance of ytterbium in the Earth crust is about 3 mg/kg.[3] The most important current (2008) sources of ytterbium are the ionic adsorption clays of southern China. The "High Yttrium" concentrate derived from some versions of these comprise about two thirds yttria by weight, and 3-4% ytterbia. As an even-numbered lanthanide, in accordance with the Oddo-Harkins rule, ytterbium is significantly more abundant than its immediate neighbors, thulium and lutetium, which occur in the same concentrate at levels of about 0.5% each. The world production of ytterbium is only about 50 tonnes per year, reflecting the fact that it finds little commercial application.[3]
Production Recovery of ytterbium from ores involves several processes which are common to most rare-earth elements: 1) processing, 2) separation of Yb from other rare earths, 3) preparation of the metal. If the starting ore is gadolinite, it is digested with hydrochloric or nitric acid which dissolves the rare-earth metals. The solution is treated with sodium oxalate or oxalic acid to precipitate rare earths as oxalates. For euxenite, or is processed either by fusion with potassium bisulfate or with hydrofluoric acid. Monazite or xenotime are heated either with sulfuric acid or with caustic soda.
5
Ytterbium Ytterbium is separated from other rare earths either by ion exchange or by reduction with sodium amalgam. In the latter method, a buffered acidic solution of trivalent rare earths is treated with molten sodium mercury alloy, which reduces and dissolves Yb3+. The alloy is treated with hydrochloric acid. The metal is extracted from the solution as oxalate and converted to oxide by heating. The oxide is reduced to metal by heating with lanthanum, aluminium, cerium or zirconium in high vacuum. The metal is purified by sublimation and collected over a condensed plate.[8]
Applications Source of gamma rays The 169Yb isotope has been used as a radiation source substitute for a portable X-ray machine when electricity was not available. Like X-rays, gamma rays pass through soft tissues of the body, but are blocked by bones and other dense materials. Thus, small 169Yb samples (which emit gamma rays) act like tiny X-ray machines useful for radiography of small objects. Experiment shows that radiographs taken with 169Yb source are roughly equivalent to those taken with X-rays having energies between 250 and 350 keV.[9]
Doping of stainless steel Ytterbium could also be used to help improve the grain refinement, strength, and other mechanical properties of stainless steel. Some ytterbium alloys have been used in dentistry.[2] [3]
Yb as dopant of active media Yb is used as dopant in optical materials, usually in the form of ions in active laser media. Several powerful double-clad fiber lasers and disk lasers use Yb3+ ions as dopant at concentration of several atomic percent. Glasses (optical fibers), crystals and ceramics with Yb3+ are used.[10] Ytterbium is often used as a doping material (as Yb3+) for high power and wavelength-tunable solid state lasers. Yb lasers commonly radiate in the 1.06–1.12 µm band being optically pumped at wavelength 900 nm–1 µm, dependently on the host and application. Small quantum defect makes Yb prospective dopant for efficient lasers and power scaling.[11] The kinetic of excitations in Yb-doped materials is simple and can be described within concept of effective cross-sections; for the most of Yb-doped laser materials (as for many other optically-pumped gain media), the McCumber relation holds,[12] [10] [13] although the application to the Yb-doped composite materials was under discussion.[14] [15] Usually, low concentrations of Yb are used. At high concentration of excitations, the Yb-doped materials show photodarkening[16] (glass fibers) or ever switch to the broadband emission [17] (crystals and ceramics) instead of the efficient laser action. This effect may be related with not only overheating, but also conditions of the charge compensation at high concentration of Yb ions.[18]
6
Ytterbium
7
Others Ytterbium metal increases its electrical resistivity when subjected to high stresses. This property is used in stress gauges to monitor ground deformations from earthquakes and explosions.[19]
Precautions Although ytterbium is fairly stable, it nevertheless should be stored in closed containers to protect it from air and moisture. All compounds of ytterbium should be treated as highly toxic although initial studies appear to indicate that the danger is limited. Ytterbium compounds are, however, known to cause skin and eye irritation and may be teratogenic.[20] Metallic ytterbium dust poses a fire and explosion hazard.[21]
See also • Erbium • Terbium • Yttrium
Further reading • Guide to the Elements – Revised Edition, Albert Stwertka, (Oxford University Press; 1998) ISBN 0-19-508083-1 • It's Elemental – Ytterbium [22]
External links • WebElements.com – Ytterbium
[23]
(also used as a reference)
References [1] M. Jackson "Magnetism of Rare Earth" The IRM quaterly col. 10, No. 3, p. 1, 2000 (http:/ / www. irm. umn. edu/ quarterly/ irmq10-3. pdf) [2] C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304814. [3] John Emsley (2003). Nature's building blocks: an A-Z guide to the elements (http:/ / books. google. co. jp/ books?id=j-Xu07p3cKwC). Oxford University Press. pp. 492–494. ISBN 0198503407. . [4] M. Jackson "Magnetism of Rare Earth" The IRM quarterly col. 10, No. 3, p. 1, 2000 (http:/ / www. irm. umn. edu/ quarterly/ irmq10-3. pdf) [5] " Chemical reactions of Ytterbium (https:/ / www. webelements. com/ ytterbium/ chemistry. html)". Webelements. . Retrieved 2009-06-06. [6] Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils; (1985). "Die Lanthanoide" (in German). Lehrbuch der Anorganischen Chemie (91–100 ed.). Walter de Gruyter. pp. 1265–1279. ISBN 3-11-007511-3. [7] James B. Hedrick. " Rare-Earth Metals (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ rare_earths/ 740798. pdf)". USGS. . Retrieved 2009-06-06. [8] Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds (http:/ / books. google. com/ books?id=Xqj-TTzkvTEC& pg=PA243). McGraw-Hill. pp. 973–975. ISBN 0070494398. . Retrieved 2009-06-06. [9] R. Halmshaw (1995). Industrial radiology: theory and practice (http:/ / books. google. co. jp/ books?id=wJqBSA1exqoC& ). Springer. pp. 168–169. ISBN 0412627809. . [10] D. Kouznetsov; J.-F. Bisson, K. Takaichi, K. Ueda (2005). " Single-mode solid-state laser with short wide unstable cavity (http:/ / josab. osa. org/ abstract. cfm?id=84730)". JOSAB 22 (8): 1605–1619. doi: 10.1364/JOSAB.22.001605 (http:/ / dx. doi. org/ 10. 1364/ JOSAB. 22. 001605). . [11] Grukh, Dmitrii A (2004). "Broadband radiation source based on an ytterbium-doped fibre with fibre-length-distributed pumping". Quantum Electronics 34: 247. doi: 10.1070/QE2004v034n03ABEH002621
Ytterbium (http:/ / dx. doi. org/ 10. 1070/ QE2004v034n03ABEH002621). [12] D. E. McCumber (1964). "Einstein relations connecting broadband emission and absorption spectra". Physical Review B 136 (4A): 954–957. doi: 10.1103/PhysRev.136.A954 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 136. A954). [13] P. C. Becker, N. A. Olson, J. R. Simpson. (1999). Erbium-doped fiber amplifiers: fundamentals and theory. Academic press. [14] D. Kouznetsov (2007). "Comment on Efficient diode-pumped Yb:Gd2SiO5 laser". Applied Physics Letters 90: 066101. doi: 10.1063/1.2435309 (http:/ / dx. doi. org/ 10. 1063/ 1. 2435309). [15] Guangjun Zhao; Liangbi Su, Jun Xu, Heping Zeng (2007). "Response to Comment on Efficient diode-pumped Yb:Gd2SiO5 laser". Applied Physics Letters 90: 066103. doi: 10.1063/1.2435314 (http:/ / dx. doi. org/ 10. 1063/ 1. 2435314). [16] Joona J. Koponen; Mikko J. Söderlund, Hanna J. Hoffman, and Simo K. T. Tammela (2006). "Measuring photodarkening from single-mode ytterbium doped silica fibers". Optics Express 14 (24): 11539–11544. doi: 10.1364/OE.14.011539 (http:/ / dx. doi. org/ 10. 1364/ OE. 14. 011539). [17] J.-F. Bisson; D. Kouznetsov, K. Ueda, S. T. Fredrich-Thornton, K. Petermann, G. Huber (2007). "Switching of emissivity and photoconductivity in highly doped Yb3+:Y2O3 and Lu2O3 ceramics". Applied Physics Letters 90: 201901. doi: 10.1063/1.2739318 (http:/ / dx. doi. org/ 10. 1063/ 1. 2739318). [18] N. V. Sochinskii; M. Abellan, J. Rodriguez-Fernandez, E. Saucedo, C. M. Ruiz, V. Bermudez (2007). "Effect of Yb concentration on the resistivity and lifetime of CdTe:Ge:Yb codoped crystals". Applied Physics Letters 91 (20): 202112. doi: 10.1063/1.2815644 (http:/ / dx. doi. org/ 10. 1063/ 1. 2815644). [19] C. K. Gupta, Nagaiyar Krishnamurthy (2004). Extractive metallurgy of rare earths (http:/ / books. google. co. jp/ books?id=F0Bte_XhzoAC& pg=PA32). CRC Press. p. 32. ISBN 0415333407. . [20] Gale, Tf (Jun 1975). "The embryotoxicity of ytterbium chloride in golden hamsters.". Teratology 11 (3): 289–95. doi: 10.1002/tera.1420110308 (http:/ / dx. doi. org/ 10. 1002/ tera. 1420110308). ISSN 0040-3709 (http:/ / worldcat. org/ issn/ 0040-3709). PMID 807987. [21] " Material safety data sheet (http:/ / www. espi-metals. com/ msds's/ ytterbium. pdf)". . Retrieved 2009-06-06. [22] http:/ / education. jlab. org/ itselemental/ ele070. html [23] http:/ / www. webelements. com/ webelements/ elements/ text/ Yb/ index. html
8
Article Sources and Contributors
Article Sources and Contributors Ytterbium Source: http://en.wikipedia.org/w/index.php?oldid=307618113 Contributors: 2over0, AWeenieMan, Addshore, Ahoerstemeier, AlimanRuna, Archimerged, Ardonik, Arkuat, Atakdoug, Barneca, Beetstra, Benbest, Benjiboi, Biruitorul, Blathnaid, BlueEarth, Bobo192, Bryan Derksen, Cacahueten, Can't sleep, clown will eat me, Chem-awb, Christian List, CommonsDelinker, Conversion script, D99figge, Darrien, David Latapie, Domitori, Donarreiskoffer, DrBob, Edgar181, El C, Elkman, Emperorbma, Encyclopedia77, Excirial, Feezo, Femto, Gene Nygaard, Gh, Grendelkhan, Hak-kâ-ngìn, Hbackman, I Like Cheeseburgers, Icairns, Ideyal, InfoCan, IvanLanin, J.delanoy, JForget, Jaraalbe, JazzWriter, Jcobb, Joanjoc, Karelj, Karlhahn, Kelovy, King of Hearts, KnowledgeOfSelf, Kralizec!, Kurykh, Kwamikagami, LA2, LarryMorseDCOhio, Lexicon, LilHelpa, Majorly, Marc Venot, Materialscientist, Mav, Mets501, MichaelBillington, Mike Rosoft, Minesweeper, Moe Epsilon, Morwen, Mouser, Mr.Z-man, Nergaal, Nihiltres, Nikai, Peruvianllama, Pharaoh of the Wizards, PinchasC, Poolkris, Pras, RPaschotta, RSido, RTC, Remember, RexNL, Reza kalani, Roberta F., Romanm, RoryReloaded, RoyBoy, Saperaud, Savidan, Schneelocke, Sengkang, Shaddack, Sl, Smilesfozwood, Snigbrook, SnowFire, Steve Hart, Stifynsemons, Stone, Syd Henderson, Tagishsimon, Tetracube, Triona, Vicki Rosenzweig, Vsmith, Warut, Whosasking, Yekrats, Yyy, 145 anonymous edits
Image Sources, Licenses and Contributors image:Yb-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Yb-TableImage.png License: GNU Free Documentation License Contributors: Cecil, CommonsDelinker, Georg Slickers, Paddy, Paginazero, Saperaud Image: Ytterbium 1-croprotated.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Ytterbium_1-croprotated.jpg License: unknown Contributors: User:Maksim, User:Materialscientist File:Euxenite - Vegusdal, Norvegia 01.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Euxenite_-_Vegusdal,_Norvegia_01.jpg License: unknown Contributors: User:Aangelo
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
9
Lutetium
1
Lutetium 71
ytterbium ← lutetium → hafnium
Y ↑
Lu ↓
Lr Periodic Table - Extended Periodic Table
General Name, Symbol, Number
lutetium, Lu, 71
Element category
transition metals
Group, Period, Block
n/a, 6, d
Appearance
silvery white
Standard atomic weight
174.967(1) g·mol
Electron configuration
[Xe] 6s 4f
−1
2
Electrons per shell
14
1
5d
2, 8, 18, 32, 9, 2 Physical properties
Phase
solid
Density (near r.t.)
9.841 g·cm−3
Liquid density at m.p.
9.3 g·cm−3
Melting point
1925 K (1652 °C, 3006 °F)
Boiling point
3675 K (3402 °C, 6156 °F)
Heat of fusion
ca. 22 kJ·mol−1
Heat of vaporization
414 kJ·mol−1
Specific heat capacity
(25 °C) 26.86 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1906
2103
2346
(2653)
(3072)
(3663)
Atomic properties
Lutetium
2
Crystal structure
hexagonal
Oxidation states
3 (weakly basic oxide)
Electronegativity
1.27 (Pauling scale)
Ionization energies (more)
1st: 523.5 kJ·mol−1 2nd: 1340 kJ·mol−1 3rd: 2022.3 kJ·mol−1
Atomic radius
174 pm
Covalent radius
187±8 pm Miscellaneous [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(r.t.) (poly) 582 nΩ·m
Thermal conductivity
(300 K) 16.4 W·m
Thermal expansion
(r.t.) (poly) 9.9 µm/(m·K)
Young's modulus
68.6 GPa
Shear modulus
27.2 GPa
Bulk modulus
47.6 GPa
−1
Poisson ratio
−1
·K
0.261
Vickers hardness
1160 MPa
Brinell hardness
893 MPa
CAS registry number
7439-94-3 Most-stable isotopes
Main article: Isotopes of lutetium iso
NA
half-life
DM
DE (MeV)
DP
173
syn
1.37 y
ε
0.671
173
174
syn
3.31 y
ε
1.374
174
1.193
176
Lu Lu
175
97.41%
175
176
2.59%
3.78×1010y
Lu Lu
Yb Yb
Lu is stable with 104 neutron β−
Hf
References
Lutetium (pronounced /ljuːˈtiːʃiəm/) is a chemical element with the symbol Lu and atomic number 71. A silvery-white rare metal, lutetium has the smallest atomic radius and is the heaviest and hardest member of the rare-earth group. One of its radioactive isotopes (176Lu) is used in nuclear technology to determine the age of meteorites. Lutetium usually occurs in association with yttrium and is sometimes used in metal alloys and as a catalyst in various chemical reactions. A strict correlation between periodic table blocks and chemical series for neutral atoms would describe lutetium as a transition metal because it is in the d-block, but it is a lanthanide according to IUPAC.[2] Lutetium is one of the rarest and is the
Lutetium most expensive rare-earth element.
Characteristics Physical Lutetium is a silvery white corrosion-resistant trivalent metal. It has the smallest atomic radius and is the heaviest and hardest of the rare earth elements.[3] Lutetium has the highest melting point of any lanthanide, probably related to the lanthanide contraction.
Chemical Lutetium metal tarnishes slowly in air and burns readily at 150 °C to form lutetium(III) oxide: 4 Lu + 3 O2 → 2 Lu2O3 Lutetium is quite electropositive and reacts slowly with cold water and quite quickly with hot water to form lutetium hydroxide: 2 Lu (s) + 6 H2O (g) → 2 Lu(OH)3 (aq) + 3 H2 (g) Lutetium metal reacts with all the halogens to form halides: 2 Lu (s) + 3 F2 (g) → 2 LuF3 (s) 2 Lu (s) + 3 Cl2 (g) → 2 LuCl3 (s) 2 Lu (s) + 3 Br2 (g) → 2 LuBr3 (s) 2 Lu (s) + 3 I2 (g) → 2 LuI3 (s) The fluoride, chloride, and bromide are white, whereas the iodide is brown. Lutetium dissolves readily in dilute sulfuric acid to form solutions containing the colorless lutetium(III) ions, which exist as a [Lu(OH2)9]3+ complex:[4] 2 Lu (s) + 3 H2SO4 (aq) → 2 Lu3+ (aq) + 3 SO2−4 (aq) + 3 H2 (g)
Compounds In all its compounds, lutetium occurs in +3 valence state. Aqueous solutions of most Lu salts are colorless and form white crystalline solids upon drying. The soluble salts, such as chloride (LuCl3), bromide (LuBr3), iodide (LuI3), nitrate, sulfate and acetate form hydrates upon crystallization. The oxide (Lu2O3), hydroxide, fluoride (LuF3), carbonate, phosphate and oxalate are insoluble in water.[5] Lutetium tantalate (LuTaO4) is the densest known stable white material (density 9.81 g/cm3)[6] and therefore is an ideal host for X-ray phosphors.[7] [8] Thoria is more dense (10 g/cm3) and is also white, but radioactive.
Isotopes Naturally occurring lutetium is composed of 1 stable isotope 175Lu (97.41% natural abundance) and 1 long-lived beta-radioactive isotope 176Lu with a half-life of 3.78×1010 years (2.59% natural abundance). The last one is used in radiometric dating (see Lutetium-hafnium dating). 33 radioisotopes have been characterized, with the most stable being naturally occurring 176Lu, and artificial isotopes 174Lu with a half-life of 3.31 years, and 173Lu with a half-life of 1.37 years. All of the remaining radioactive isotopes have half-lives that are less than 9 days, and the majority of these have half-lives that are less
3
Lutetium than half an hour. This element also has 18 meta states, with the most stable being (T½=160.4 days), 174mLu (T½=142 days) and 178mLu (T½=23.1 minutes).
4 177m
Lu
The known isotopes of lutetium range in atomic weight from 149.973 (150Lu) to 183.961 (184Lu). The primary decay mode before the most abundant stable isotope, 175Lu, is electron capture (with some alpha and positron emission), and the primary mode after is beta emission. The primary decay products before 175Lu are element 70 (ytterbium) isotopes and the primary products after are element 72 (hafnium) isotopes.
History Lutetium (Latin Lutetia meaning Paris) was independently discovered in 1907 by French scientist Georges Urbain,[9] Austrian mineralogist Baron Carl Auer von Welsbach, and American chemist Charles James. [10] All of these men found lutetium as an impurity in the mineral ytterbia which was thought by Swiss chemist Jean Charles Galissard de Marignac (and most others) to consist entirely of the element ytterbium. The separation of lutetium from Marignac's ytterbium was first described by Urbain and the naming honor therefore went to him. He chose the names neoytterbium (new ytterbium) and lutecium for the new element but neoytterbium was eventually reverted back to ytterbium and in 1949 the spelling of element 71 was changed to lutetium. The dispute on the priority of the discovery is documented in two articles in which Urbain and von Welsbach accuse each other of publishing results influenced by the published research of the other.[11] [12] The Commission on Atomic Mass, which was responsible for the attribution of the names for the new elements, settled the dispute in 1909 by granting priority to Urbain and adopting his names as official ones. An obvious problem with this decision was that Urbain was one of the four members of the commission.[13] Welsbach proposed the names cassiopium for element 71 (after the constellation Cassiopeia) and aldebaranium for the new name of ytterbium but these naming proposals were rejected (although many German scientists in the 1950s called the element 71 cassiopium). The irony of all this is that Charles James, who had modestly stayed out of the argument as to priority, worked on a much larger scale than the others, and undoubtedly possessed the largest supply of lutetium at the time.[14]
Occurrence and production
Lutetium
Found with almost all other rare-earth metals but never by itself, lutetium is very difficult to separate from other elements. The principal commercially viable ore of lutetium is the rare earth phosphate mineral monazite: (Ce, La, etc.) PO4 which contains 0.003% of the element. The abundance of lutetium in the Earth crust is only about 0.5 mg/kg. The main mining areas are China, United States, Brazil, India, Sri Lanka and Australia. The world production of lutetium (in the form of oxide) is about 10 tonnes per year.[14] Pure lutetium metal has only relatively recently been isolated and is Monazite very difficult to prepare. It is one of the rarest and most expensive of the rare earth metals with the price about US$ 10,000 per kg.[15] [16] Crushed minerals are treated with hot concentrated sulfuric acid to produce water-soluble sulfates of rare earths. Thorium precipitates out of solution as hydroxide and is removed. After that the solution is treated with ammonium oxalate to convert rare earths in to their insoluble oxalates. The oxalates are converted to oxides by annealing. The oxides are dissolved in nitric acid that excludes one of the main components, cerium, whose oxide is insoluble in HNO3. Several rare earth metals, including Lu, are separated as a double salt with ammonium nitrate by crystallization. Lutetium is separated by ion exchange. In this process, rare-earth ions are sorbed onto suitable ion-exchange resin by exchange with hydrogen, ammonium or cupric ions present in the resin. Lutetium salts are then selectively washed out by suitable complexing agent. Lutetium metal is then obtained by reduction of anhydrous LuCl3 or LuF3 by either an alkali metal or alkaline earth metal.[5] 2 LuCl3 → 2 Lu + 3 CaCl2
Applications Because of the rarity and high price, lutetium has very few commercial uses. However, stable lutetium can be used as catalysts in petroleum cracking in refineries and can also be used in alkylation, hydrogenation, and polymerization applications. Some other applications include: • Lutetium-176 (176Lu) has been used to date the age of meteorites.[17] • Lutetium aluminium garnet (Al5Lu3O12) has been proposed for use as a lens material in high refractive index immersion lithography.[18] • Lutetium-177 (177Lu), when bound to Octreotate (a somatostatin analogue), is used experimentally in targeted radionuclide therapy for neuroendocrine tumours.[19] • Cerium-doped lutetium oxyorthosilicate (LSO) is currently the preferred compound for detectors in positron emission tomography (PET.) [20] [21] • Use as a pure beta emitter, using lutetium which has been exposed to neutron activation. A tiny amount of lutetium is added as a dopant to gadolinium gallium garnet (GGG), which is used in magnetic bubble memory devices.[22]
5
Lutetium
Precautions Like other rare-earth metals, lutetium is regarded as having a low degree of toxicity but it and especially its compounds should be handled with care nonetheless. Metal dust of this element is a fire and explosion hazard. Lutetium plays no biological role in the human body.
References [1] Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf). CRC press. 2000. ISBN 0849304814. . [2] " IUPAC Provisional Recommendations for the Nomenclature of Inorganic Chemistry (online draft of an updated version of the "Red Book" IR 3-6) (http:/ / www. iupac. org/ reports/ provisional/ abstract04/ connelly_310804. html)". 2004. . Retrieved 2009-06-06. [3] Parker, Sybil P. (1984). Dictionary of Scientific and Technical Terms, 3rd ed. New York: McGraw-Hill. [4] " Chemical reactions of Lutetium (https:/ / www. webelements. com/ lutetium/ chemistry. html)". Webelements. . Retrieved 2009-06-06. [5] Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds (http:/ / books. google. com/ books?id=Xqj-TTzkvTEC& pg=PA243). McGraw-Hill. p. 510. ISBN 0070494398. . Retrieved 2009-06-06. [6] Blasse, G (1994). "Luminescence of materials based on LuTaO4". Journal of Alloys and Compounds 209: 1. doi: 10.1016/0925-8388(94)91069-3 (http:/ / dx. doi. org/ 10. 1016/ 0925-8388(94)91069-3). edit (http:/ / en. wikipedia. org/ wiki/ Template:cite_doi/ 10. 1016. 2f0925-8388. 2894. 2991069-3) [7] Shigeo Shionoya (1998). Phosphor handbook (http:/ / books. google. co. jp/ books?id=lWlcJEDukRIC& pg=PA846). CRC Press. p. 846. ISBN 0849375606. . [8] C. K. Gupta, Nagaiyar Krishnamurthy (2004). Extractive metallurgy of rare earths (http:/ / books. google. co. jp/ books?id=F0Bte_XhzoAC& pg=PA32). CRC Press. p. 32. ISBN 0415333407. . [9] M. G. Urbain (1908). " Un nouvel élément, le lutécium, résultant du dédoublement de l'ytterbium de Marignac (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k3099v/ f759. table)". Comptes rendus 145: 759–762. . [10] " Separation of Rare Earth Elements (http:/ / acswebcontent. acs. org/ landmarks/ landmarks/ rareearth/ discovery. html)". . [11] C. Auer v. Welsbach (1908). "Die Zerlegung des Ytterbiums in seine Elemente". Monatshefte für Chemie 29 (2): 181–225. doi: 10.1007/BF01558944 (http:/ / dx. doi. org/ 10. 1007/ BF01558944). [12] G. Urbain (1909). "Lutetium und Neoytterbium oder Cassiopeium und Aldebaranium -- Erwiderung auf den Artikel des Herrn Auer v. Welsbach". Monatshefte für Chemie 31 (10): I. doi: 10.1007/BF01530262 (http:/ / dx. doi. org/ 10. 1007/ BF01530262). [13] F. W. Clarke, W. Ostwald, T. E. Thorpe, G. Urbain (1909). "Bericht des Internationalen Atomgewichts-Ausschusses für 1909". Berichte der deutschen chemischen Gesellschaft 42 (1): 11–17. doi: 10.1002/cber.19090420104 (http:/ / dx. doi. org/ 10. 1002/ cber. 19090420104). [14] John Emsley (2001). Nature's building blocks: an A-Z guide to the elements (http:/ / books. google. com/ books?id=Yhi5X7OwuGkC& pg=PA241). US: Oxford University Press. pp. 240–242. ISBN 0198503415. . [15] James B. Hedrick. " Rare-Earth Metals (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ rare_earths/ 740798. pdf)". USGS. . Retrieved 2009-06-06. [16] Stephen B. Castor and James B. Hedrick. " Rare Earth Elements (http:/ / www. rareelementresources. com/ i/ pdf/ RareEarths-CastorHedrickIMAR7. pdf)". . Retrieved 2009-06-06. [17] Muriel Gargaud, Hervé Martin, Philippe Claeys (2007). Lectures in Astrobiology (http:/ / books. google. com/ books?id=3uYmP0K5PXEC& pg=PA52). Springer. p. 51. ISBN 3540336923. . [18] Yayi Wei, Robert L. Brainard (2009). Advanced Processes for 193-NM Immersion Lithography (http:/ / books. google. com/ books?id=Sx39H8XR1FcC& pg=PA12). SPIE Press. p. 12. ISBN 0819475572. . [19] Helmut Sigel (2004). Metal complexes in tumor diagnosis and as anticancer agents (http:/ / books. google. com/ books?id=ZtRdbUNbPn8C& pg=PA98). CRC Press. p. 98. ISBN 0824754948. . [20] Wahl RL (2002). "Instrumentation". Principles and Practice of Positron Emission Tomography. Philadelphia: Lippincott: Williams and Wilkins. p. 51. [21] Daghighian, F. Shenderov, P. Pentlow, K.S. Graham, M.C. Eshaghian, B. Melcher, C.L. Schweitzer, J.S. (1993). "Evaluation of cerium doped lutetium oxyorthosilicate (LSO)scintillation crystals for PET". Nuclear Science 40 (4): 1045–1047. doi: 10.1109/23.256710 (http:/ / dx. doi. org/ 10. 1109/ 23. 256710). [22] J. W. Nielsen, S. L. Blank, D. H. Smith, G. P. Vella-Coleiro, F. B. Hagedorn, R. L. Barns and W. A. Biolsi (1974). "Three garnet compositions for bubble domain memories". Journal of Electronic Materials 3 (3): 693–707. doi: 10.1007/BF02655293 (http:/ / dx. doi. org/ 10. 1007/ BF02655293).
6
Lutetium • Guide to the Elements - Revised Edition, Albert Stwertka, (Oxford University Press; 1998) ISBN 0-19-508083-1
External links • WebElements.com - Lutetium (http:/ / www. webelements. com/ webelements/ elements/ text/ Lu/ index. html) (also used as a reference) • It's Elemental - Lutetium (http:/ / education. jlab. org/ itselemental/ ele071. html) • pure Lutetium >99,9% picture in the element collection from Heinrich Pniok (http:/ / www. pse-mendelejew. de/ bilder/ lu. jpg) • Lutetium (http:/ / www. mrteverett. com/ Chemistry/ pdictable/ q_elements. asp?Symbol=Lu) pnb:میتیتل
7
Article Sources and Contributors
Article Sources and Contributors Lutetium Source: http://en.wikipedia.org/w/index.php?oldid=308692924 Contributors: AWeishaupt, Abtinb, Ahoerstemeier, Alchemist-hp, AlimanRuna, Archimerged, Arkuat, Bakabaka, Beetstra, Benbest, Bencherlite, Big Bird, BlueEarth, Borislav Dopudja, Bryan Derksen, Bucephalus, Cacycle, CanisRufus, Canthusus, Conversion script, DMacks, David Latapie, Donarreiskoffer, EL Willy, Edgar181, Element16, Emperorbma, Epolk, Eras-mus, Eu-151, Femto, Flying Jazz, Fredrik, Giftlite, Greatpatton, Grendelkhan, Gurch, Hadal, Hak-kâ-ngìn, Halfdan, Harland1, Hdossa, Icairns, Ideyal, JForget, Jaraalbe, Joanjoc, KDerbyshire, Karelj, Karnesky, Kelovy, Kwamikagami, Kyoko, LA2, LarryMorseDCOhio, LilHelpa, Marc Venot, Materialscientist, Mav, Minesweeper, Mrfb061, Nergaal, Nihiltres, PeepP, PeterJeremy, Physchim62, Pleasantville, Plexust, Poolkris, Pras, Psyche825, RTC, Remember, Reyk, Riana, Roberta F., Romanm, RoyBoy, SajmonDK, Saperaud, Schneelocke, Sengkang, Sergio.ballestrero, Sfuerst, Silver Spoon, Sl, Slicky, Squids and Chips, Steve Hart, Stifynsemons, Stone, Tagishsimon, Tetracube, Thingg, V1adis1av, Vlad4599, Vsmith, Walkerma, Warut, Woohookitty, Wrenchelle, Yekrats, Yyy, Ziggy Sawdust, 85 anonymous edits
Image Sources, Licenses and Contributors image:Lu-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Lu-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Paddy, Saperaud Image: Lutetium amp.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Lutetium_amp.jpg License: Public Domain Contributors: Original uploader was Tomihahndorf at de.wikipedia Image:Monazit - Mosambik, O-Afrika.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Monazit_-_Mosambik,_O-Afrika.jpg License: unknown Contributors: User:Ra'ike
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
8
Hafnium
1
Hafnium 72
lutetium ← hafnium → tantalum
Zr ↑
Hf ↓
Rf Periodic Table - Extended Periodic Table
General Name, Symbol, Number
hafnium, Hf, 72
Element category
transition metals
Group, Period, Block
4, 6, d
Appearance
steel grey
Standard atomic weight
178.49(2) g·mol
Electron configuration
[Xe] 4f
−1
14
Electrons per shell
2
2
5d 6s
2, 8, 18, 32, 10, 2 Physical properties
Phase
solid
Density (near r.t.)
13.31 g·cm−3
Liquid density at m.p.
12 g·cm−3
Melting point
2506 K (2233 °C, 4051 °F)
Boiling point
4876 K (4603 °C, 8317 °F)
Heat of fusion
27.2 kJ·mol−1
Heat of vaporization
571 kJ·mol−1
Specific heat capacity
(25 °C) 25.73 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
2689
2954
3277
3679
4194
4876
Atomic properties
Hafnium
2
Crystal structure
hexagonal
Oxidation states
4, 3, 2 (amphoteric oxide)
Electronegativity
1.3 (Pauling scale)
Ionization energies (more)
1st: 658.5 kJ·mol−1 2nd: 1440 kJ·mol−1 3rd: 2250 kJ·mol−1
Atomic radius
159 pm
Covalent radius
175±10 pm Miscellaneous [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(20 °C) 331 n Ω·m
Thermal conductivity
(300 K) 23.0 W·m
Thermal expansion
(25 °C) 5.9 µm·m
Speed of sound (thin rod)
(20 °C) 3010 m/s
Young's modulus
78 GPa
Shear modulus
30 GPa
Bulk modulus
110 GPa
Poisson ratio
0.37
Mohs hardness
5.5
Vickers hardness
1760 MPa
Brinell hardness
1700 MPa
CAS registry number
−1
−1
−1
−1
·K
·K
7440-58-6 Most-stable isotopes
Main article: Isotopes of hafnium iso 172
Hf
NA syn
half-life
0.350
172
α
2.495
170
2.446
178
0.373
182
2×1015 y
176
5.206%
176
177
18.606%
177
178
27.297%
178
Hf Hf
178m2
Hf
syn
Hf is stable with 106 neutron
31 y 179
180
35.1%
180
182
Hf
syn
Yb
Hf is stable with 105 neutron
13.629%
Hf
Lu
Hf is stable with 104 neutron
179
Hf
DP
ε
0.162%
Hf
DE (MeV)
1.87 y
174
Hf
DM
IT
Hf
Hf is stable with 107 neutron Hf is stable with 108 neutron
9×106 y
β−
Ta
Hafnium
3 References
Hafnium (pronounced /ˈhæfniəm/) is a chemical element with the symbol Hf and atomic number 72. A lustrous, silvery gray, tetravalent transition metal, hafnium chemically resembles zirconium and is found in zirconium minerals. Its existence was predicted by Dmitri Mendeleev in 1869. Hafnium was the second-to-last element of those with stable isotopes to be discovered. It was found by Dirk Coster and Georg von Hevesy in 1923 in Copenhagen, Denmark, and named Hafnia after the Latin name for "Copenhagen". Hafnium is used in filaments, electrodes, and some semiconductor fabrication processes for integrated circuits at 45 nm and smaller feature lengths. Its large neutron capture cross-section makes hafnium a good material for neutron absorption in control rods in nuclear power plants. Some superalloys used for special applications contain hafnium in combination with niobium, titanium, or tungsten.
History In his report on The Periodic Law of the Chemical Elements, in 1869, Dmitri Mendeleev had implicitly predicted the existence of a heavier analog of titanium and zirconium. At the time of his formulation in 1871, Mendeleev believed that the elements were ordered by their atomic masses and placed lanthanum (element 57) in the spot below zirconium. The exact placement of the elements and the location of missing elements was done by determining the specific weight of the elements and comparing the chemical and physical properties.[2] The X-ray spectroscopy done by Henry Moseley in 1914 The hafnium seal of the Faculty of Science of the University of Copenhagen
showed a direct dependency between spectral line and effective nuclear charge. This led to the nuclear charge, or atomic number of an element, being used to ascertain its place within the periodic table. With this method, Moseley determined the number of lanthanoids and showed the gaps in the atomic number sequence at numbers 43, 61, 72, and 75.[3] The discovery of the gaps led to an extensive search for the missing elements. In 1914, several people claimed the discovery after Henry Moseley predicted the gap in the periodic table for the then-undiscovered element 72.[4] Georges Urbain asserted that he found element 72 in the rare earth elements in 1907 and published his results on celtium in 1911.[5] Neither the spectra nor the chemical behavior matched with the element found later, and therefore his claim was turned down after a long standing controversy.[6] The controversy was partly due to the fact that the chemists favored the chemical techniques which lead to the discovery of celtium, while the physicists relied on the use of the new X-ray spectroscopy method that proved that the substances discovered by Urbain did not contain element 72.[6] By early 1923, several physicists and chemists such as Niels Bohr[7] and Charles R. Bury[8] suggested that element 72 should resemble zirconium and therefore was not part of the rare earth elements group. These suggestions were based on Bohr's theories of the atom, the X-ray spectroscopy of Mosley, and the chemical arguments of Friedrich Paneth.[9]
Hafnium
4
Encouraged by these suggestions and by the reappearance in 1922 of Urbain's claims that element 72 was a rare earth element discovered in 1911, Dirk Coster and Georg von Hevesy were motivated to search for the new element in zirconium ores.[10] Hafnium was discovered by the two in 1923 in Copenhagen, Denmark, validating the original 1869 prediction of Mendeleev.[11] [12] It was ultimately found in zircon in Norway through X-ray spectroscopy analysis.[13] The place where the discovery took place led to the element being named for the Latin name for "Copenhagen", Hafnia, the home town of Niels Bohr. [14] Today, the Faculty of Science of the University of Copenhagen uses in its seal a stylized image of hafnium.[15] Hafnium was separated from zirconium through repeated recrystallization of the double ammonium or potassium fluorides by Valdemar Thal Jantzen and von Hevesey.[16] Anton Eduard van Arkel and Jan Hendrik de Boer were the first prepare metallic hafnium by passing hafnium tetra-iodide vapor over a heated tungsten filament in 1924.[17] [18] This process for differential purification of zirconium and hafnium is still in use today.[19] In 1923, four predicted elements were still missing from the periodic table: 43 (technetium) and 61 (promethium) are radioactive elements and are only present in trace amounts in the environment,[20] thus making elements 75 (rhenium) and 72 (hafnium) the last two unknown non-radioactive elements. Since rhenium was discovered in 1925,[21] hafnium was the next to last element with stable isotopes to be discovered.
Characteristics Hafnium is a shiny, silvery, ductile metal that is corrosion-resistant and chemically similar to [19] zirconium. The physical properties of hafnium metal samples are markedly affected by zirconium impurities, as these two elements are among the most difficult ones to separate because of their chemical A hafnium crystal bar, made using the crystal bar process similarity.[19] A notable physical difference between them is their density (zirconium being about half as dense as hafnium). The most notable physical property of hafnium is its high thermal neutron-capture cross-section, and the nuclei of several hafnium isotopes can each absorb multiple neutrons.[19] Hafnium does react in air to form a protective film that prevents any further reaction.
Isotopes At least 34 isotopes of hafnium have been observed, ranging in mass number from 153 to 186.[22] [23] The five stable isotopes are in the range of 176 to 180. The radioactive isotopes' half-lifes range from only 400 ms for 153Hf,[23] to 2.0 petayears (1015 years) for the most stable one, 174Hf.[22] The nuclear isomer 178m2Hf is also a source of cascades of gamma rays whose energies total 2.45 MeV per decay.[24] It is notable because it has the highest excitation energy of any comparably long-lived isomer of any element. One gram of this pure isotope could release approximately 1330 megajoules of energy, the equivalent of exploding about 317 kilograms (700 pounds) of TNT. Possible applications requiring such highly concentrated energy storage are of interest. For example, it has been studied as a possible
Hafnium
5
power source for gamma ray lasers.[25]
Chemistry As a tetravalent transition metal, hafnium forms various inorganic compounds, generally in the oxidation state of +4. The metal is resistant to concentrated alkalis, but halogens react with it to form hafnium tetrahalides.[26] At higher temperatures, hafnium reacts with oxygen, nitrogen, carbon, boron, sulfur, and silicon.[26] Due to the lanthanide contraction of the elements in the fifth period, zirconium and hafnium have nearly identical ionic radii. The ionic radius of Zr4+ is 0.79 Ångström and that of Hf4+ is 0.78 Ångström.[26] This similarity results in nearly identical chemical behavior and in the formation of similar chemical compounds.[26] The chemistry of hafnium is so similar to that of zirconium that a separation on chemical reactions was not possible, only the physical properties of the compounds differ. The melting points and boiling points of the compounds and the solubility in solvents are the major differences in the chemistry of these twin elements.[27] Hafnium dioxide
Like zirconium, hafnium reacts with halogens forming the tetrahalogen compound with the oxidation state of +4 for hafnium. Hafnium(IV) chloride and hafnium(IV) iodide have some applications in the production and purification of hafnium.[27] The white hafnium oxide (HfO2), with a melting point of 2812 °C and a boiling point of roughly 5100 °C, is very similar to zirconia, but slightly basic.[27] Hafnium carbide is the most refractory binary compound known, with a melting point over 3890 °C, and hafnium nitride is the most refractory of all known metal nitrides, with a melting point of 3310 °C.[26] This has led to proposals that hafnium or its carbides might be useful as construction materials that are subjected to very high temperatures. The mixed carbide tantalum hafnium carbide (Ta4HfC5) possesses the highest melting point of any currently known compound, 4215 °C.[28]
Occurrence Hafnium is estimated to make up about 5.8 ppm of the Earth's upper crust by weight. It does not exist as a free element in nature, but is found combined in solid solution for zirconium in natural zirconium compounds such as zircon, ZrSiO4, which usually has a about 1 - 4 % of the Zr replaced by Hf. Rarely, the Hf/Zr ratio increases during crystallization to give the isostructural mineral 'hafnon' (Hf,Zr)SiO4, with atomic Hf > Zr.[29] An old (obsolete) name for a variety of zircon containing unusually high Hf content is alvite [30]
Zircon crystal from Tocantins, Brazil (unknown scale)
Hafnium
6
A major source of zircon (and hence hafnium) ores are heavy mineral sands ore deposits, pegmatites particularly in Brazil and Malawi, and carbonatite intrusions particularly the Crown Polymetallic Deposit at Mount Weld, Western Australia. A potential source of hafnium is trachyte tuffs containing rare zircon-hafnium silicates eudialyte or armstrongite, [31] at Dubbo in New South Wales, Australia.
Production The heavy mineral sands ore deposits of the titanium ores ilmenite and rutile yield most of the mined zirconium and therefore also most the hafnium.[32] Zirconium is a good fuel-rod cladding metal, with the desirable properties of a very low neutron capture cross-section and good chemical stability at high temperatures. However, because of hafnium's neutron-absorbing properties, hafnium impurities in zirconium would cause it to be far less useful for nuclear reactor applications. Thus a nearly complete separation of zirconium and hafnium is necessary for their use in nuclear power. The production of hafnium free zirconium is the main source for hafnium.[19] Several details contribute to fact that there are only a few technical uses for hafnium. First, the close similarity between hafnium and zirconium makes it possible to use zirconium for most of the applications. Second, hafnium was first available as pure metal after the use in the nuclear industry for hafnium-free zirconium in the late 1950s. Furthermore, the low abundance and the difficult separation techniques necessary make it a scarce commodity.[19] Hafnium and zirconium have nearly identical chemistry,
A lump of hafnium which has been
which makes the two difficult to separate.[33] The first oxidized on one side and exhibits thin film optical effects. used methods of fractionated crystallization of [16] ammonium fluoride salts or the fractionated [17] distillation of the chloride were not suitable for an industrial scale production. After zirconium was chosen as material for the nuclear reactor program in the 1940s, a separation method had to be developed. Liquid-liquid extraction processes with a wide variety of solvents were developed and are still used for the production of hafnium.[34] About half of all hafnium metal manufactured is produced as a by-product of zirconium refinement. The end product of the separation is hafnium(IV) chloride.[35] The conversion to the metal is done through reducing hafnium(IV) chloride with magnesium or sodium in the Kroll process.[36] HfCl4 + 2 Mg (1100 °C) → 2 MgCl2 + Hf Further purification is done by a chemical transport reaction developed by Arkel and de Boer. In a closed vessel, hafnium reacts with iodine at temperatures of 500 °C forming hafnium(IV) iodide; at a tungsten filament of 1700 °C the reverse reaction happens and the iodine and hafnium are set free. The hafnium forms a solid coating at the tungsten filament and the iodine can react with additional hafnium resulting in a steady turn over.[18] [27] Hf + 2 I2 (500 °C) → HfI4 HfI4 (1700 °C) → Hf + 2 I2
Hafnium
7
Applications Most of the hafnium produced is used in the production of control rod for nuclear reactors.[34]
Nuclear reactors The nuclei of several hafnium isotopes can each absorb multiple neutrons. This makes hafnium a good material for use in the control rods for nuclear reactors. Its neutron-capture cross-section is about 600 times that of zirconium. (Other elements that are good neutron-absorbers for control rods are cadmium and boron.) Excellent mechanical properties and exceptional corrosion-resistance properties allow its use in the harsh environment of a pressurized water reactors.[34] The German research reactor FRM II uses hafnium as a neutron absorber.[37]
Alloys Hafnium is used in iron, titanium, niobium, tantalum, and other metal alloys. An alloy used for liquid rocket thruster nozzles, for example the main engine of the Apollo Lunar Modules is C103, which consists of 89% niobium, 10% hafnium and 1% titanium.[38] Small additions of hafnium increase the adherence of protective oxide scales on nickel based alloys. It improves thereby the corrosion resistance especially under cyclic temperature conditions that tend to break oxide scales by inducing thermal stresses between the bulk material and the oxide layer.[39] [40] [41] Hafnium containing rocket nozzle of the Apollo Lunar Module in the lower right corner
Microprocessors The electronics industry discovered that hafnium-based compound can be employed in gate insulators in the 45 nm generation of integrated circuits from Intel, IBM and others.[42] [43] Hafnium oxide-based compounds are practical high-k dielectrics, allowing reduction of the gate leakage current which improves performance at such scales.[44] [45]
Other uses Due to its heat resistance and its affinity to oxygen and nitrogen, hafnium is a good scavenger for oxygen and nitrogen in gas-filled and incandescent lamps. Hafnium is also used as the electrode in plasma cutting because of its ability to shed electrons into air,[46] The high energy content of 178m2Hf is the concern of a DARPA funded program in the US. This program should determine the possibility of using a nuclear isomer of hafnium (the above mentioned 178m2Hf) to construct high yield weapons with X-ray triggering mechanisms—an application of induced gamma emission. That work follows over two decades of basic research by an international community[47] into the means for releasing
Hafnium
8
the stored energy upon demand. There is considerable opposition to this program[48] because uninvolved countries might perceive an "isomer weapon gap" that would justify their further development and stockpiling of nuclear weapons. A related proposal is to use the same isomer to power Unmanned Aerial Vehicles,[49] which could remain airborne for months at a time.
Precautions Care needs to be taken when machining hafnium because, like its sister metal zirconium, when hafnium is divided into fine particles, it is pyrophoric and can ignite spontaneously in air—similar to that obtained in Dragon's Breath. Compounds that contain this metal are rarely encountered by most people. The pure metal is not considered toxic, but hafnium compounds should be handled as if they were toxic because the ionic forms of metals are normally at greatest risk for toxicity, and limited animal testing has been done for hafnium compounds.[50]
Dragon's Breath at night
See also • Nuclear isomer • Induced gamma emission • Zircon
External links • Hafnium [51] at Los Alamos National Laboratory's periodic table of the elements • Hafnium Technical & Safety Data [53] • NLM Hazardous Substances Databank – Hafnium, elemental • Intel Shifts from Silicon to Lift Chip Performance [55] • Hafnium-based Intel 45nm Process Technology [56]
[52]
[54]
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] Kaji, Masanori (2002). " D. I. Mendeleev's concept of chemical elements and The Principles of Chemistry (http:/ / www. scs. uiuc. edu/ ~mainzv/ HIST/ awards/ OPA Papers/ 2005-Kaji. pdf)" (pdf). Bulletin for the History of Chemistry 27: 4. . Retrieved 2008-08-20. [3] Heilbron, John L. (1966). "The Work of H. G. J. Moseley". Isis 57: 336. doi: 10.1086/350143 (http:/ / dx. doi. org/ 10. 1086/ 350143). [4] Heimann, P. M. (1967). "Moseley and celtium: The search for a missing element". Annals of Science 23: 249. doi: 10.1080/00033796700203306 (http:/ / dx. doi. org/ 10. 1080/ 00033796700203306). [5] Urbain, M. G. (1911). " Sur un nouvel élément qui accompagne le lutécium et le scandium dans les terres de la gadolinite: le celtium (On a new element that accompanies lutetium and scandium in gadolinite: celtium) (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k3105c/ f141. table)" (in French). Comptes rendus: 141. . Retrieved 2008-09-10. [6] Mel'nikov, V. P. (1982). "Some Details in the Prehistory of the Discovery of Element 72". Centaurus 26: 317. doi: 10.1111/j.1600-0498.1982.tb00667.x (http:/ / dx. doi. org/ 10. 1111/ j. 1600-0498. 1982. tb00667. x).
Hafnium [7] Bohr, Niels. The Theory of Spectra and Atomic Constitution: Three Essays (http:/ / ia311508. us. archive. org/ 0/ items/ TheTheoryOfSpectraAndAtomicConstitution/ HTM/ 00000131. htm). p. 114. . [8] Bury, Charles R. (1921). "Langmuir's Theory of the Arrangement of Electrons in Atoms and Molecules". J. Amer. Chem. Soc. 43: 1602. doi: 10.1021/ja01440a023 (http:/ / dx. doi. org/ 10. 1021/ ja01440a023). [9] Paneth, F. A. (1922). "Das periodische System (The periodic system)" (in German). Ergebnisse der Exakten Naturwissenschaften 1. p. 362. [10] Urbain, M. G. (1922). " Sur les séries L du lutécium et de l'ytterbium et sur l'identification d'un celtium avec l'élément de nombre atomique 72 (The L series from luthetium to ytterbium and the identification of element 72 celtium (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k3127j/ f1348. table)" (in French). Comptes rendus 174: 1347. . Retrieved 2008-10-30. [11] Coster, D.; Hevesy, G. (1923). "On the Missing Element of Atomic Number 72". Nature 111: 79. doi: 10.1038/111079a0 (http:/ / dx. doi. org/ 10. 1038/ 111079a0). [12] Hevesy, G. (1925). "The Discovery and Properties of Hafnium". Chemical Reviews 2: 1. doi: 10.1021/cr60005a001 (http:/ / dx. doi. org/ 10. 1021/ cr60005a001). [13] von Hevesy, Georg (1923). "Über die Auffindung des Hafniums und den gegenwärtigen Stand unserer Kenntnisse von diesem Element". Berichte der deutschen chemischen Gesellschaft (A and B Series) 56: 1503. doi: 10.1002/cber.19230560702 (http:/ / dx. doi. org/ 10. 1002/ cber. 19230560702). [14] Scerri, Eric R. (1994). "Prediction of the nature of hafnium from chemistry, Bohr's theory and quantum theory". Annals of Science 51: 137. doi: 10.1080/00033799400200161 (http:/ / dx. doi. org/ 10. 1080/ 00033799400200161). [15] " University Life 2005 (http:/ / www. ku. dk/ publikationer/ pdf/ arkiv/ aarsberetning/ University_life_2005. pdf)" (pdf). University of Copenghagen. p. 43. . Retrieved 2008-11-02. [16] van Arkel, A. E.; de Boer, J. H. (1924). "Die Trennung von Zirkonium und Hafnium durch Kristallisation ihrer Ammoniumdoppelfluoride (The separation of zirconium and hafnium by crystallization of the double ammonium fluorides)" (in German). Zeitschrift für anorganische und allgemeine Chemie 141: 284. doi: 10.1002/zaac.19241410117 (http:/ / dx. doi. org/ 10. 1002/ zaac. 19241410117). [17] van Arkel, A. E.; de Boer, J. H. (1924). "Die Trennung des Zirkoniums von anderen Metallen, einschließlich Hafnium, durch fraktionierte Distillation (The separation of zirconium and hafnium by fractionated distillation)" (in German). Zeitschrift für anorganische und allgemeine Chemie 141: 289. doi: 10.1002/zaac.19241410118 (http:/ / dx. doi. org/ 10. 1002/ zaac. 19241410118). [18] van Arkel, A. E.; de Boer, J. H. (1925). "Darstellung von reinem Titanium-, Zirkonium-, Hafnium- und Thoriummetall (Production of pure titanium, zirconium, hafnium and Thorium metal)" (in German). Zeitschrift für anorganische und allgemeine Chemie 148: 345. doi: 10.1002/zaac.19251480133 (http:/ / dx. doi. org/ 10. 1002/ zaac. 19251480133). [19] Schemel, J. H. (1977). ASTM Manual on Zirconium and Hafnium (http:/ / books. google. com/ books?id=dI_LssydVeYC). ASTM International. pp. 1-5. ISBN 9780803105058. . [20] Curtis, David; Fabryka-Martin, June; Dixon, Pauland; Cramer, Jan (1999). "Nature’s uncommon elements: plutonium and technetium". Geochimica et Cosmochimica Acta 63: 275. doi: 10.1016/S0016-7037(98)00282-8 (http:/ / dx. doi. org/ 10. 1016/ S0016-7037(98)00282-8). [21] Noddack, W.; Tacke, I.; Berg, O (1925). "Die Ekamangane". Naturwissenschaften 13: 567. doi: 10.1007/BF01558746 (http:/ / dx. doi. org/ 10. 1007/ BF01558746). [22] EnvironmentalChemistry.com. " Hafnium Nuclides / Isotopes (http:/ / environmentalchemistry. com/ yogi/ periodic/ Hf-pg2. html#Nuclides)". Periodic Table of Elements. J.K. Barbalace. . Retrieved 2008-09-10. [23] Georges, Audi (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [24] " WWW Table of Radioactive Isotopes (http:/ / ie. lbl. gov/ toi/ nuclide. asp?iZA=720778)". Lawrence Berkeley National Laboratory Isotopes Project and Lund University. . Retrieved 2008-09-10. [25] Collins, C. B.; Zoita, N. C.; Davanloo, F.; Yoda, Y.; Uruga, T.; Pouvesle, J. M.; Popescu, I. I. (2004). "Nuclear resonance spectroscopy of the 31-yr isomer of Hf-178". Laser Physics Letters 2: 162. doi: 10.1002/lapl.200410154 (http:/ / dx. doi. org/ 10. 1002/ lapl. 200410154). [26] " Los Alamos National Laboratory – Hafnium (http:/ / periodic. lanl. gov/ elements/ 72. html)". . Retrieved 2008-09-10. [27] Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils; (1985) (in German). Lehrbuch der Anorganischen Chemie (91-100 ed.). Walter de Gruyter. pp. 1056–1057. ISBN 3110075113. [28] Deadmore, D. L. (1964). " Vaporization of Tantalum-Carbide-Hafnium-Carbide Solid Solutions at 2500 to 3000 K (http:/ / ntrs. nasa. gov/ archive/ nasa/ casi. ntrs. nasa. gov/ 19650001401_1965001401. pdf)" (PDF). NASA. . Retrieved 2008-11-02.
9
Hafnium [29] Deer, William Alexander; Howie, R.A.; Zussmann, J. (1982). The Rock-Forming Minerals, volume 1A: Orthosilicates. Longman Group Limited. pp. 418-442. ISBN 0582465265. [30] Lee, O. Ivan (1928). "The Mineralogy of Hafnium" (pdf). Chemical Reviews 5: 17. doi: 10.1021/cr60017a002 (http:/ / dx. doi. org/ 10. 1021/ cr60017a002). [31] " Dubbo Zirconia Project Fact Sheet (http:/ / www. alkane. com. au/ projects/ nsw/ dubbo/ DZP Summary June07. pdf)" (pdf). Alkane Resources Limited. June 2007. . Retrieved 2008-09-10. [32] Gambogi, Joseph. " Yearbook 2008: Zirconium and Hafnium (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ zirconium/ myb1-2007-zirco. pdf)" (pdf). . Retrieved 2008-10-27. [33] Larsen, Edwin; Fernelius W., Conard; Quill, Laurence (1943). "Concentration of Hafnium. Preparation of Hafnium-Free Zirconia". Ind. Eng. Chem. Anal. Ed. 15: 512. doi: 10.1021/i560120a015 (http:/ / dx. doi. org/ 10. 1021/ i560120a015). [34] Hedrick, James B.. " Hafnium (http:/ / minerals. er. usgs. gov/ minerals/ pubs/ commodity/ zirconium/ 731798. pdf)" (pdf). United States Geological Survey. . Retrieved 2008-09-10. [35] Griffith, Robert F. (1952). " Zirconium and hafnium (http:/ / digicoll. library. wisc. edu/ cgi-bin/ EcoNatRes/ EcoNatRes-idx?type=turn& entity=EcoNatRes. MinYB1952v1. p1172& isize=M)". Minerals yearbook metals and minerals (except fuels). The first production plants Bureau of Mines. p. 1162–1171. . [36] Gilbert, H. L.; Barr, M. M. (1955). "Preliminary Investigation of Hafnium Metal by the Kroll Process". Journal of the Electrochemical Society 102: 243. doi: 10.1149/1.2430037 (http:/ / dx. doi. org/ 10. 1149/ 1. 2430037). [37] " Forschungsreaktor München II (FRM-II): Standort und Sicherheitskonzept (http:/ / www. ssk. de/ werke/ volltext/ 1995/ ssk9512. pdf)" (pdf). Strahlenschutzkommission. 1996-02-07. . Retrieved 2008-09-22. [38] Hebda, John (2001). " Niobium alloys and high Temperature Applications (http:/ / www. cbmm. com. br/ portug/ sources/ techlib/ science_techno/ table_content/ sub_3/ images/ pdfs/ 016. pdf)" (pdf). CBMM. . Retrieved 2008-09-04. [39] Maslenkov, S. B.; Burova, N. N.; Khangulov, V. V. (1980). "Effect of hafnium on the structure and properties of nickel alloys". Metal Science and Heat Treatment 22: 283. doi: 10.1007/BF00779883 (http:/ / dx. doi. org/ 10. 1007/ BF00779883). [40] Beglov, V. M.; Pisarev, B. K.; Reznikova, G. G. (1992). "Effect of boron and hafnium on the corrosion resistance of high-temperature nickel alloys". Metal Science and Heat Treatment 34: 251. doi: 10.1007/BF00702544 (http:/ / dx. doi. org/ 10. 1007/ BF00702544). [41] Voitovich, R. F.; Golovko, É. I. (1975). "Oxidation of hafnium alloys with nickel". Metal Science and Heat Treatment 17: 207. doi: 10.1007/BF00663680 (http:/ / dx. doi. org/ 10. 1007/ BF00663680). [42] US patent 6013553 (http:/ / v3. espacenet. com/ textdoc?DB=EPODOC& IDX=US6013553) [43] Markoff, John (2007-01-27). " Intel Says Chips Will Run Faster, Using Less Power (http:/ / www. nytimes. com/ 2007/ 01/ 27/ technology/ 27chip. html)". New York Times. . Retrieved 2008-09-10. [44] Fulton, III, Scott M. (January 27, 2007). " Intel Reinvents the Transistor (http:/ / www. betanews. com/ article/ Intel_Reinvents_the_Transistor/ 1169872301)". BetaNews. . Retrieved 2007-01-27. [45] Robertson, Jordan (January 27, 2007). " Intel, IBM reveal transistor overhaul (http:/ / www. washingtonpost. com/ wp-dyn/ content/ article/ 2007/ 01/ 27/ AR2007012700152. html)". The Associated Press. . Retrieved 2008-09-10. [46] Ramakrishnany, S.; Rogozinski, M. W. (1997). " Properties of electric arc plasma for metal cutting (http:/ / www. iop. org/ EJ/ article/ 0022-3727/ 30/ 4/ 019/ d70419. pdf)" (pdf). Journal of Physics D: Applied Physics 30: 636. doi: 10.1088/0022-3727/30/4/019 (http:/ / dx. doi. org/ 10. 1088/ 0022-3727/ 30/ 4/ 019). . [47] " Isomer Triggering History, (http:/ / www. hafniumisomer. org/ isomer/ IGEhistory. htm)". The Center for Quantum Electronics, The University of Texas at Dallas. . Retrieved 2008-09-10. [48] Schwarzschild, Bertram (May 2004). "Conflicting Results on a Long-Lived Nuclear Isomer of Hafnium Have Wider Implications". Physics Today 57: 21. doi: 10.1063/1.1768663 (http:/ / dx. doi. org/ 10. 1063/ 1. 1768663). [49] Graham-Rowe, Duncan (2003-02-19). " Nuclear-powered drone aircraft on drawing board (http:/ / www. newscientist. com/ article/ dn3406-nuclearpowered-drone-aircraft-on-drawing-board. html)". New Scientist. . Retrieved 2008-06-06. [50] " Occupational Safety & Health Administration: Hafnium (http:/ / www. osha. gov/ SLTC/ healthguidelines/ hafnium/ index. html)". U.S. Department of Labor. . Retrieved 2008-09-10. [51] [52] [53] [54] [55]
http:/ / periodic. lanl. gov/ elements/ 72. html http:/ / periodic. lanl. gov/ default. htm http:/ / www. americanelements. com/ hf. htm http:/ / toxnet. nlm. nih. gov/ cgi-bin/ sis/ search/ r?dbs+ hsdb:@term+ @na+ @rel+ hafnium,+ elemental http:/ / online. wsj. com/ article/ SB119481053795589302. html
[56] http:/ / www. intel. com/ technology/ 45nm/ index. htm?iid=homepage+ marquee_45nm
10
Article Sources and Contributors
Article Sources and Contributors Hafnium Source: http://en.wikipedia.org/w/index.php?oldid=306215581 Contributors: Aastrup, Ahoerstemeier, Alchemist-hp, AlimanRuna, Archimerged, ArielGold, Arkuat, Ashleychamberpot, Athaler, Auximines, AxelHarvey, Axiosaurus, B07, Baccyak4H, Beetstra, Benbest, Berek Halfhand, BillFlis, Blaxthos, BlueEarth, Bobblewik, Bobo192, Borislav Dopudja, Brockert, Bryan Derksen, Bwieliczko, C0nanPayne, CYD, CanisRufus, Carnildo, CharlesC, ChicXulub, Chowbok, Chris 73, ChrisUTMBA, Conversion script, Darth Panda, David Latapie, David elliott lewis, Deglr6328, Deli nk, Delta G, Derek.cashman, Dfrg.msc, Discospinster, Drac2000, Edgar181, El C, Emperorbma, Encyclopedia77, Epbr123, Erik Zachte, Evil saltine, Femto, Fuzzy22, GMU22LamarButler, Gazimoff, Gene Nygaard, GeoGreg, Giftlite, Gilliam, Gpkh, Greatpatton, Greg L, Grendelkhan, Gtg204y, Hak-kâ-ngìn, HazyM, Helixblue, Hermann Luyken, Hqb, Iammaxus, Icairns, Ideyal, Ilikepie2221, Jaraalbe, Jasongagich, JdH, Kelovy, Kortaggio, Krellis, Ksheka, Kubigula, Kwamikagami, LA2, LarryMorseDCOhio, Lawnrrd, Lewis R, LorenzoB, Magnus Manske, Marc Venot, Materialscientist, MatthieuV, Mattisse, Maury Markowitz, Mav, Maxironchin, Maxis ftw, Minesweeper, Mortdefides, Motor, NawlinWiki, Ndenison, Nergaal, Neurolysis, Nihiltres, NrDg, PMLawrence, Paul Drye, Peregrine981, Physchim62, Plexust, Poolkris, Potatoswatter, Pras, Protonk, Qdr, RCWord, RTC, Ragingbullfrog, Remember, Reza kalani, Rholton, Rjwilmsi, Robert Merkel, Roberta F., Rolinator, Romanm, Ryan Postlethwaite, S, Sam Hocevar, Saperaud, Sbharris, Schneelocke, Seandop, SebastianHelm, Sengkang, Sfuerst, Shadanan, Shaddack, Shipmaster, Shizaru, Signalhead, Simesa, Simn19, Sl, Slakr, Slowbie, Smokefoot, Speculatrix, SpookyMulder, Stifynsemons, Stone, Student7, Tagishsimon, The Lizard Wizard, Thingg, Thortveitite, Timeastor, Tubedogg, User27091, Vikingforties, VinnieCool, Vsmith, Vuo, Warut, Watch37264, Welsh, Whitepaw, Wikieed, Yakushima, Yekrats, Yyy, 212 anonymous edits
Image Sources, Licenses and Contributors image:Hf-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Hf-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Conscious, Paddy, Saperaud Image: Hafnium bits.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Hafnium_bits.jpg License: unknown Contributors: Original uploader was Deglr6328 at en.wikipedia Image:Sigilum Facultatis Naturalis.JPG Source: http://en.wikipedia.org/w/index.php?title=File:Sigilum_Facultatis_Naturalis.JPG License: logo Contributors: Aastrup, MBisanz Image:Hf-crystal bar.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Hf-crystal_bar.jpg License: unknown Contributors: User:Alchemist-hp Image:Hafnium(IV) oxide.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Hafnium(IV)_oxide.jpg License: Public Domain Contributors: User:Walkerma Image:Zircão.jpeg Source: http://en.wikipedia.org/w/index.php?title=File:Zircão.jpeg License: Creative Commons Attribution-Sharealike 2.0 Contributors: Zimbres Image:Hafnium lump thin film effects.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Hafnium_lump_thin_film_effects.jpg License: GNU Free Documentation License Contributors: Deglr6328 Image:Apollo AS11-40-5866.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Apollo_AS11-40-5866.jpg License: Public Domain Contributors: Neil A. Armstrong Image:Dragonsbreath.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Dragonsbreath.jpg License: GNU Free Documentation License Contributors: Kyle Sternberg
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11
Tantalum
1
Tantalum 73
hafnium ← tantalum → tungsten
Nb ↑
Ta ↓
Db Periodic Table - Extended Periodic Table
General Name, Symbol, Number
tantalum, Ta, 73
Element category
transition metals
Group, Period, Block
5, 6, d
Appearance
gray blue
Standard atomic weight
180.94788(2) g·mol
Electron configuration
[Xe] 4f
−1
14
Electrons per shell
3
2
5d 6s
2, 8, 18, 32, 11, 2 Physical properties
Phase
solid
Density (near r.t.)
16.69 g·cm−3
Liquid density at m.p.
15 g·cm−3
Melting point
3290 K (3017 °C, 5463 °F)
Boiling point
5731 K (5458 °C, 9856 °F)
Heat of fusion
36.57 kJ·mol−1
Heat of vaporization
732.8 kJ·mol−1
Specific heat capacity
(25 °C) 25.36 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
3297
3597
3957
4395
4939
5634
Atomic properties
Tantalum
2
Crystal structure
body centered cubic
Oxidation states
5, 4, 3, 2, -1 (mildly acidic oxide)
Electronegativity Ionization energies
1.5 (Pauling scale) 1st: 761 kJ/mol 2nd: 1500 kJ/mol
Atomic radius
146 pm
Covalent radius
170±8 pm Miscellaneous [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(20 °C) 131 n Ω·m
Thermal conductivity
(300 K) 57.5 W·m
Thermal expansion
(25 °C) 6.3 µm·m
Speed of sound (thin rod)
(20 °C) 3400 m/s
Young's modulus
186 GPa
Shear modulus
69 GPa
Bulk modulus
200 GPa
Poisson ratio
0.34
Mohs hardness
6.5
Vickers hardness
873 MPa
Brinell hardness
800 MPa
CAS registry number
7440-25-7 Most-stable isotopes
−1
−1
−1
−1
·K
·K
Tantalum
3
Main article: Isotopes of tantalum iso
NA
half-life
DM
DE (MeV)
DP
177
syn
56.56 h
ε
1.166
177
178
syn
2.36 h
ε
1.910
178
179
syn
1.82 a
ε
0.110
179
180
syn
8.125 h
ε
0.854
180
β−
0.708
180
ε
0.929
180
β
−
0.783
180
IT
0.075
180
−
1.814
182
−
1.070
183
Ta Ta Ta Ta
180m
Ta
181
Ta
0.012%
99.988%
>1.2×1015 y (not observed)
Hf Hf W Hf W Ta
Ta is stable with 108 neutron
syn
114.43 d
β
183
syn
5.1 d
β
Ta
Hf
181
182
Ta
Hf
W W
References
Tantalum (pronounced /ˈtæntələm/) (formerly tantalium /tænˈtæliəm/) is a chemical element with the symbol Ta and atomic number 73. A rare, hard, blue-gray, lustrous transition metal, tantalum is highly corrosion resistant and occurs naturally in the mineral tantalite, always together with the chemically similar niobium. It is part of the refractory metals group, which are widely used as minor component in alloys. The chemical inertness of tantalum makes it a valuable substance for laboratory equipment and a substitute for platinum, but its main use today is in tantalum capacitors.
History Tantalum was discovered in Sweden in 1802 by Anders Ekeberg. One year earlier, Charles Hatchett had discovered the element columbium.[2] In 1809, the English chemist William Hyde Wollaston compared the oxides derived from both columbium—columbite, with a density 5.918 g/cm3, and tantalum—tantalite, with a density 7.935 g/cm3, and concluded that the two oxides, despite the significant difference in density, were identical; thus he kept the name tantalum.[3] After Friedrich Wöhler confirmed these results it was believed that columbium and tantalum were the same element. This conclusion was disputed in 1846 by the German chemist Heinrich Rose, who argued that there were two additional elements in the tantalite sample, and named them after children of Tantalus: niobium (from Niobe, the goddess of tears), and pelopium (from Pelops).[4] [5] The element pelopium later was identified as a mixture of tantalum and niobium, while the niobium was identical to the columbium already discovered in 1801 by Hattchet. The differences between tantalum and niobium were unequivocally demonstrated in 1864 by Christian Wilhelm Blomstrand,[6] and Henri Etienne Sainte-Claire Deville, as well as Louis J. Troost, who determined the formulas of some of the compounds in 1865[6] [7] and
Tantalum finally by the Swiss chemist Jean Charles Galissard de Marignac,[8] in 1866, who all proved that there were only two elements. These discoveries did not stop scientists from publishing articles about ilmenium until 1871.[9] De Marignac was the first to prepare the metal in 1864, when he reduced tantalum chloride by heating it in an atmosphere of hydrogen.[10] Early investigators were only able to isolate impure metal and the first relatively pure ductile metal was produced by Werner von Bolton in 1903. Wires made with tantalum metal were used for light bulbs until tungsten replaced it.[11] Its name is derived from the character Tantalus, father of Niobe in Greek mythology, who was punished after death by being condemned to stand knee-deep in water with perfect fruit growing above his head, both of which eternally tantalized him - if he bent to drink the water, it drained below the level he could reach, and if he reached for the fruit, the branches moved out of his grasp.[12] Ekeberg wrote "This metal I call tantalum … partly in allusion to its incapacity, when immersed in acid, to absorb any and be saturated."[13] For many years, the commercial technology for separating tantalum from niobium involved the fractional crystallization of potassium heptafluorotantalate away from potassium oxypentafluoroniobate monohydrate, a process discovered by Jean Charles Galissard de Marignac in 1866. The method has been supplanted by solvent extraction from fluoride-containing solutions.[7] The mining of coltan, a tantalum ore, in the conflict regions of Democratic Republic of the Congo raised ethical questions and human rights issues, and endangered wildlife.[14] [15] [16]
Characteristics Physical Tantalum is dark, dense, ductile, very hard, easily fabricated, and highly conductive of heat and electricity. The metal is renowned for its resistance to corrosion by acids; in fact, at temperatures below 150 °C tantalum is almost completely immune to attack by the normally aggressive aqua regia. It can be dissolved with hydrofluoric acid or acidic solutions containing the fluoride ion and sulfur trioxide, as well as with a solution of potassium hydroxide. Tantalum's high melting point of 3017 °C (boiling point 5458 °C) is exceeded only by tungsten and rhenium for metals, and carbon.
Chemical It is able to form oxides with the oxidation states +5 (Ta2O5) and +4 (TaO2),[17] The most stable oxidation state is +5, tantalum pentoxide.[17] Tantalum pentoxide is the starting material for several tantalum compounds. The compounds are created by dissolving the pentoxide in basic hydroxide solutions or by melting it in another metal oxide. Such examples are lithium tantalate (LiTaO3) and lanthanum tantalate (LaTaO4). In the lithium tantalate, the tantalate ion TaO3 is not alone, but part of a perovskite-like structure; while the lanthanum niobate contains lone TaO4 ions.[17] The fluorides of tantalum can be used for its separation from niobium.[18] Tantalum forms halogen compounds in the oxidation states of +5, +4, and +3 of the type TaX5, TaX4, and TaX3, although multi core complexes and substoichiometric compounds are also known.[17] [19] Tantalum pentafluoride (TaF5) is a white solid with a melting point of 97.0 °C and tantalum pentachloride (TaCl5) is a white solid with a melting point of 247.4 °C. Tantalum
4
Tantalum pentachloride is hydrolyzed by water and reacts with additional tantalum at elevated temperatures by forming the black and highly hygroscopic tantalum tetrachloride (TaCl4). While the trihalogen compounds can be obtained by reduction of the pentahalogenes with hydrogen, the dihalogen compounds do not exist.[17] A tantalum-tellurium alloy forms [17] quasicrystals. Tantalum compounds with oxidation states as low as -1 have been [20] reported in 2008. Like most of the other refractory metals, the hard forms are stable nitrides and carbides. Tantalum carbide, like the more commonly used tungsten carbide, is a very hard ceramic used in cutting tools. Tantalum (III) nitride is used as a thin film insulator in some microelectronic fabrication processes.[21] Los Alamos National Laboratory scientists have developed a tantalum carbide-graphite composite material that is one of the hardest materials ever synthesized. Korean researchers have developed an amorphous tantalum-tungsten-copper alloy which is more flexible and two to three times stronger than traditional steel alloys.[22] There are two tantalum aluminides, TaAl3 and Ta3Al; they are stable, refractory and reflective, and have been proposed[23] as mirror coatings for use in the IR.
Isotopes Natural tantalum consists of two isotopes: 180mTa (0.012%) and 181Ta (99.988%). 181Ta is a stable isotope. 180mTa (m denotes a metastable state) is predicted to decay in three ways: isomeric transition to the ground state of 180Ta, beta decay to 180W, or electron capture to 180 Hf. However, any radioactivity of this nuclear isomer was never observed. Only a lower limit on its half life of over 1015 years has been set. The ground state of 180Ta has a half life of only 8 hours. 180mTa is the only naturally occurring nuclear isomer (excluding radiogenic and cosmogenic short-living nuclides). It is also the rarest isotope in the Universe, taking into account the elemental abundance of tantalum and isotopic abundance of 180mTa in the natural mixture of isotopes (and again excluding radiogenic and cosmogenic short-living nuclides).[24] Tantalum has been proposed as a "salting" material for nuclear weapons (cobalt is another, better-known salting material). A jacket of 181Ta, irradiated by the intense high-energy neutron flux from an exploding thermonuclear weapon, would transmute into the radioactive isotope 182Ta with a half-life of 114.43 days and produce approximately 1.12 MeV of gamma radiation, significantly increasing the radioactivity of the weapon's fallout for several months. Such a weapon is not known to have ever been built, tested, or used.[25]
5
Tantalum
Occurrence Tantalum is estimated to make up about 1 ppm[26] or 2 ppm[19] of the Earth's crust by weight. There are many species of tantalum minerals, only some of which are so far being used by industry as raw materials: tantalite, microlite, wodginite, euxenite, polycrase. Tantalite (Fe,Mn) Ta2O6 is the most important mineral for tantalum extraction. Tantalite has the same mineral structure as columbite (Fe,Mn) (Ta,Nb)2O6; when there is more Ta than Nb it is called tantalite and when there is more Nb than Ta is it called columbite (or niobite). The high density of tantalite and other tantalum Tantalite, Pilbarra district, Australia containing minerals makes the use of gravitational separation the best method. Other minerals include samarskite and fergusonite. The main production of tantalum occurs in Australia, where the largest producer, Talison Minerals (formerly part of the Sons of Gwalia company), operates the Wodgina mine. The mine produces tantalum oxide from tantalite.[27] While the large scale production of niobium in Brazil and Canada, yields also a comparable small amount of tantalum, other countries like China, Ethiopia and Mozambique mine the minerals with a higher rate of tantalum and produce a significant amount of the world production. Tantalum is also produced in Thailand and Malaysia as a by-product of tin mining and smelting. During gravity separation of the ore from placer deposits not only yield Cassiterite (SnO2) but also small amounts of tantalite are enriched in the final ore concentrated. The tin smelter slag derived from ore of these deposits contains significant amounts of tantalum and is leached from the slag.[7] [28] Future large sources of supply, in order of magnitude, are being explored in Saudi Arabia, Egypt, Greenland, China, Mozambique, Canada, Australia, the United States, Finland and Brazil.[29] [30] In central Africa the colloquial term coltan is used to refer to the two minerals equally, an example being the Democratic Republic of the Congo which the United States Geological Survey reports in its 2006 yearbook as having produced a little less than 1% of the world's tantalum for the past four years.[28] Ethical questions have been raised about responsible corporate behavior, human rights and endangered wildlife, due to the exploitation of resources such as coltan in the conflict regions of the Congo.[31] According to United Nations report[32] smuggling and exportation of coltan helped fuel the war in the Congo, a crisis that has resulted in approximately 5.4 million[33] deaths since 1998 – making it the world’s deadliest documented conflict since World War II.
Production Several steps are involved in the extraction of tantalum from tantalite. First the mineral is crushed and concentrated by gravity separation. This is generally carried out near the mine site. Further processing by chemical separation is usually done by treating the ores with a mixture of hydrofluoric acid and sulfuric acid at over 90°C. This causes the tantalum and niobium to dissolve as complex fluorides and be separated from the impurities. Ta2O5 + 14 HF → 2 H2[TaF7] + 5 H2O Nb2O5 + 10 HF → 2 H2[NbOF5] + 3 H2O
6
Tantalum
7
The first industrial scale separation developed by de Marignac used the difference in solubility between the complex niobium and tantalum fluorides K2[NbOF5]•H2O (dipotassium oxypentafluoroniobate monohydrate) and K2[TaF7] (dipotassium heptafluorotantalate) in water. Newer processes use the liquid extraction of the fluorides from aqueous solution by organic solvents such as cyclohexanone.[18] The complex niobium and tantalum fluorides are extracted separately from the organic solvent with water and either precipitated by the addition of potassium fluoride to produce a potassium fluoride complex, or precipitated with ammonia as the pentoxide:[17] H2[TaF7] + KF → K2[TaF7]↓ + HF 2 H2[TaF7] + 14 NH4OH → Ta2O5↓ + 14 NH4F + 9 H2O The resulting potassium fluorotantalate salt is generally treated by reduction with molten sodium to produce a coarse tantalum powder.[34]
Applications Electronics The major use for tantalum, as the metal powder, is in the production of electronic components, mainly capacitors and some high-power resistors[35] . Tantalum electrolytic capacitors exploit the tendency of tantalum to form a protective oxide surface layer, using tantalum powder, pressed into a pellet shape, as one "plate" of the capacitor, the oxide as the dielectric, and an electrolytic solution or conductive solid as the other "plate". Because the dielectric layer can be very thin (thinner than the similar layer in, for instance, an aluminium electrolytic capacitor), a high capacitance can be achieved in a small volume. Because of the size and weight advantages, tantalum capacitors are attractive for portable telephones, personal computers, and automotive electronics.[36] Tantalum electrolytic capacitor
Alloys Tantalum is also used to produce a variety of alloys that have high melting points, are strong and have good ductility. Alloyed with other metals, it is also used in making carbide tools for metalworking equipment and in the production of superalloys for jet engine components, chemical process equipment, nuclear reactors, and missile parts.[36] [37] Because of its ductility, tantalum can be drawn into fine wires or filaments, which are used for evaporating metals such as aluminium. Due to the fact that it resists attack by body fluids and is nonirritating, tantalum is widely used in making surgical instruments and implants. For example, porous tantalum coatings are used in the construction of orthopedic implants due to tantalum's ability to form a direct bond to hard tissue.[38]
Tantalum
8
Tantalum is inert against most acids except hydrofluoric acid and hot sulfuric acid, also hot alkaline solutions cause tantalum to corrode. This property makes it an ideal metal for chemical reaction vessels and pipes for corrosive liquids. Heat exchanging coils for the steam heating of hydrochloric acid are made from tantalum.[39] Tantalum was extensively used in the production of ultra high frequency electron tubes for radio transmitters. The tantalum is capable of capturing oxygen and nitrogen by forming nitrides and oxides and therefore helps to sustain the high vacuum needed for the tubes.[18] [39]
Other uses The oxide is used to make special high refractive index glass for camera lenses.[40] The high melting point and oxidation resistance lead to the use of the metal in the production of vacuum furnace parts. Due to its high density, shaped charge and explosively formed penetrator liners have been constructed from tantalum.[41] Tantalum greatly increases the armor penetration capabilities of a shaped charge due to its high density and high melting point.[42] [43] It is also occasionally used in precious watches e.g. from Hublot, Montblanc and Panerai.
Precautions Compounds containing tantalum are rarely encountered in the laboratory. The metal is highly biocompatible and is used for body implants and coatings, therefore attention may be focused on other elements or the physical nature of the chemical compound.[44] A single study[45] is the only reference in literature ever linking tantalum to local sarcomas. It is possible the result was due to other factors not considered in the study. The study was quoted in IARC Monograph vol. 74 which includes the following "Note to the reader": "Inclusion of an agent in the Monographs does not imply that it is a carcinogen, only that the published data have been examined."[46]
External links • WebElements.com: Tantalum
[47]
• Tantalum-Niobium International Study Center
[48]
pnb:ملاٹنیٹ
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] Griffith, William P.; Morris, Peter J. T. (2003). " Charles Hatchett FRS (1765-1847), Chemist and Discoverer of Niobium (http:/ / www. jstor. org/ stable/ 3557720)". Notes and Records of the Royal Society of London 57 (3): 299. doi: 10.1098/rsnr.2003.0216 (http:/ / dx. doi. org/ 10. 1098/ rsnr. 2003. 0216). . [3] Wollaston, William Hyde (1809). " On the Identity of Columbium and Tantalum (http:/ / www. jstor. org/ stable/ 107264)". Philosophical Transactions of the Royal Society of London 99: 246–252. doi: 10.1098/rstl.1809.0017 (http:/ / dx. doi. org/ 10. 1098/ rstl. 1809. 0017). . [4] Rose, Heinrich (1844). " Ueber die Zusammensetzung der Tantalite und ein im Tantalite von Baiern enthaltenes neues Metall (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k15148n/ f327. table)" (in German). Annalen der Physik 139 (10): 317–341. doi: 10.1002/andp.18441391006 (http:/ / dx. doi. org/ 10. 1002/ andp. 18441391006). . [5] Rose, Heinrich (1847). " Ueber die Säure im Columbit von Nordamérika (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k15155x/ f586. table)" (in German). Annalen der Physik 146 (4): 572–577. doi: 10.1002/andp.18471460410 (http:/ / dx. doi. org/ 10. 1002/ andp. 18471460410). .
Tantalum [6] Marignac, Blomstrand, H. Deville, L. Troost und R. Hermann (1866). "Tantalsäure, Niobsäure, (Ilmensäure) und Titansäure". Fresenius' Journal of Analytical Chemistry 5 (1): 384–389. doi: 10.1007/BF01302537 (http:/ / dx. doi. org/ 10. 1007/ BF01302537). [7] Gupta, C. K.; Suri, A. K. (1994). Extractive Metallurgy of Niobium. CRC Press. ISBN 0849360714. [8] Marignac, M. C. (1866). " Recherches sur les combinaisons du niobium (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k34818t/ f4. table)" (in French). Annales de chimie et de physique 4 (8): 7–75. . [9] Hermann, R. (1871). "Fortgesetzte Untersuchungen über die Verbindungen von Ilmenium und Niobium, sowie über die Zusammensetzung der Niobmineralien (Further research about the compounds of ilmenium and niobium, as well as the composition of niobium minerals)" (in German). Journal für Praktische Chemie 3 (1): 373–427. doi: 10.1002/prac.18710030137 (http:/ / dx. doi. org/ 10. 1002/ prac. 18710030137). [10] " Niobium (http:/ / nautilus. fis. uc. pt/ st2. 5/ scenes-e/ elem/ e04100. html)". Universidade de Coimbra. . Retrieved 2008-09-05. [11] Bowers, B. (2001). "Scanning Our Past from London The Filament Lamp and New Materials". Proceedings of the IEEE 89 (3): 413. doi: 10.1109/5.915382 (http:/ / dx. doi. org/ 10. 1109/ 5. 915382). [12] Aycan, Mugla, Sule (2005). "Chemistry Education and Mythology". Journal of Social Sciences 1 (4): 238–239. [13] Greenwood, Norman N.; Earnshaw, A. (1997), Chemistry of the Elements (2nd ed.), Oxford: Butterworth-Heinemann, p. 1138, ISBN 0-7506-3365-4 [14] " Congo's Bloody Coltan (http:/ / www. pulitzercenter. org/ openitem. cfm?id=177)". . Retrieved 2009-08-08. [15] " Congo War and the Role of Coltan (http:/ / www1. american. edu/ ted/ ice/ congo-coltan. htm)". . Retrieved 2009-08-08. [16] " Coltan mining in the Congo River Basin (http:/ / www. panda. org/ what_we_do/ where_we_work/ congo_basin_forests/ problems/ mining/ coltan_mining/ )". . Retrieved 2009-08-08. [17] Holleman, A. F., Wiberg, E., Wiberg, N. (2007). Lehrbuch der Anorganischen Chemie, 102nd ed.. de Gruyter. ISBN 978-3-11-017770-1. [18] Soisson, Donald J.; McLafferty, J. J.; Pierret, James A. (1961). "Staff-Industry Collaborative Report: Tantalum and Niobium". Ind. Eng. Chem. 53 (11): 861–868. doi: 10.1021/ie50623a016 (http:/ / dx. doi. org/ 10. 1021/ ie50623a016). [19] Agulyansky, Anatoly (2004). The Chemistry of Tantalum and Niobium Fluoride Compounds (http:/ / books. google. de/ books?id=Z-4QXNB5Hp8C). Elsevier. ISBN 9780444516046. . Retrieved 2008-09-02. [20] doi: 10.1021/om701189e (http:/ / dx. doi. org/ 10. 1021/ om701189e) [21] Tsukimoto, S.; Moriyama, M.; Murakami, Masanori (1961). "Microstructure of amorphous tantalum nitride thin films". Thin Solid Films 460 (1-2): 222–226. doi: 10.1016/j.tsf.2004.01.073 (http:/ / dx. doi. org/ 10. 1016/ j. tsf. 2004. 01. 073). [22] Arirang, TV (2005-05-06). " Researchers Develop New Alloy (http:/ / english. chosun. com/ w21data/ html/ news/ 200505/ 200505060005. html)". Digital Chosunilbo (English Edition) : Daily News in English About Korea. . Retrieved 2008-12-22. [23] " US Patent 5923464 - Substance for front surface mirror (http:/ / www. patentstorm. us/ patents/ 5923464/ description. html)". . Retrieved 2008-12-22. [24] Georges, Audi (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3–128. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [25] Win, David Tin; Masum, Al (2003) (PDF). Weapons of Mass Destruction (http:/ / www. journal. au. edu/ au_techno/ 2003/ apr2003/ aujt6-4_article07. pdf). 6. pp. 199–219. . [26] Emsley, John (2001). "Tantalum". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. p. 420. ISBN 0198503407. [27] " Wodgina Operations (http:/ / www. talison. com. au/ operations. html)". Talison Minerals. 2008. . Retrieved 2009-07-31. [28] Papp, john F. (2006). " 2006 Minerals Yearbook Nb & Ta (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ niobium/ #pubs)". US Geological Survey. . Retrieved 2008-06-03. [29] " Tantalum supplement (http:/ / www. noventa. net/ pdf/ presentations/ tanatalumSCR_presentation. pdf)" (PDF). Mining Journal. 2007-November. . Retrieved 2008-06-03. [30] " International tantalum resources — exploration and mining (http:/ / www. doir. wa. gov. au/ documents/ gswa/ gsdMRB_22_chap10. pdf)" (pdf). GSWA Mineral Resources Bulletin 22 (10). . [31] Hayes, Karen; Burge, Richard. Coltan Mining in the Democratic Republic of Congo: How tantalum-using industries can commit to the reconstruction of the DRC. 1–64. ISBN 1903703107. [32] " S/2003/1027 (http:/ / www. un. org/ Docs/ journal/ asp/ ws. asp?m=S/ 2003/ 1027)". 2003-10-26. . Retrieved 2008-04-19. [33] " Special Report: Congo (http:/ / www. theirc. org/ special-report/ congo-forgotten-crisis. html)". International Rescue Committee. . Retrieved 2008-04-19.
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Tantalum [34] " Extraction/refining (http:/ / tanb. org/ tantalum)". T.I.C.. . Retrieved 2009-07-07. [35] " What is a resistor? (http:/ / www. wisegeek. com/ what-is-a-resistor. htm)". . Retrieved 2009-08-08. [36] " Commodity Report 2008: Tantalum (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ niobium/ mcs-2008-tanta. pdf)" (PDF). United States Geological Survey. . Retrieved 2008-10-24. [37] Buckman Jr., R. W. (2000). "New applications for tantalum and tantalum alloys". JOM Journal of the Minerals, Metals and Materials Society 52 (3): 40. doi: 10.1007/s11837-000-0100-6 (http:/ / dx. doi. org/ 10. 1007/ s11837-000-0100-6). [38] Cohen, R. (2006). "Applications of porous tantalum in total hip arthroplasty". Journal of the American Academy of Orthopaedic Surgeons 14: 646. [39] Balke, Clarence W.. "Columbium and Tantalum". Industrial and Engineering Chemistry 20 (10): 1166. [40] Musikant, Solomon (1985). " Optical Glas Composition (http:/ / books. google. de/ books?id=iJEXMF3JBtQC& pg=PA28)". Optical Materials: An Introduction to Selection and Application. CRC Press. p. 28. ISBN 9780824773090. . [41] Nemat-Nasser, Sia; Isaacs, Jon B.; Liu, Mingqi (1998). "Microstructure of high-strain, high-strain-rate deformed tantalum". Acta Materialia 46: 1307. doi: 10.1016/S1359-6454(97)00746-5 (http:/ / dx. doi. org/ 10. 1016/ S1359-6454(97)00746-5). [42] Walters, William; Cooch, William; Burkins, Matthew (2001). "The penetration resistance of a titanium alloy against jets from tantalum shaped charge liners". International Journal of Impact Engineering 26: 823. doi: 10.1016/S0734-743X(01)00135-X (http:/ / dx. doi. org/ 10. 1016/ S0734-743X(01)00135-X). [43] Russell, Alan M.; Lee, Kok Loong (2005). Structure-property relations in nonferrous metals (http:/ / books. google. de/ books?id=fIu58uZTE-gC& pg=PA129& lpg=PP128#PPA218). Hoboken, NJ: Wiley-Interscience. p. 218. ISBN 9780471649526. . [44] Matsuno H, Yokoyama A, Watari F, Uo M, Kawasaki T. (2001). " Biocompatibility and osteogenesis of refractory metal implants, titanium, hafnium, niobium, tantalum and rhenium. Biocompatibility of tantalum (http:/ / www. ncbi. nlm. nih. gov/ pubmed/ 11336297)". Biomaterials 22: 1253. doi: 10.1016/S0142-9612(00)00275-1 (http:/ / dx. doi. org/ 10. 1016/ S0142-9612(00)00275-1). . [45] Oppenheimer, B.S.; Oppenheimer, E.T.; Danishefsky, I.; Stout, A.P. (1956). "Carcinogenic effects of metals in rodent". Cancer Research 16: 439. [46] " Surgical implants and other foreign bodies (http:/ / www. inchem. org/ documents/ iarc/ vol74/ implants. html)". IARC. 1999. . Retrieved 2009-06-03. [47] http:/ / www. webelements. com/ tantalum/ [48] http:/ / tanb. org/
10
Article Sources and Contributors
Article Sources and Contributors Tantalum Source: http://en.wikipedia.org/w/index.php?oldid=308850832 Contributors: Agentbla, Ahoerstemeier, [email protected], AlimanRuna, Anclation, Archimerged, Atmoz, Avant Guard, Avono, Axeman89, Bdiscoe, Benbest, BlueEarth, Bobathon71, Borislav Dopudja, Bryan Derksen, CHAD3112, CYD, Carnildo, CharlesC, Cimbalom, Conversion script, Cyrius, DMacks, DVD R W, DanielCristofani, Darrien, Dave6, David Latapie, Deglr6328, Dina, Dismas, Doc Tropics, Donarreiskoffer, Donko XI, Dori, DrBob, Dynzmoar, EPO, Edgar181, El C, Emperorbma, Epbr123, Erik Zachte, FT2, Femto, Fivemack, GJeffery, Gene Nygaard, GlobeGores, Greatpatton, Grendelkhan, Gwernol, Hak-kâ-ngìn, Hallmm, Helge Skjeveland, Herbee, Heron, HorsePunchKid, IanGM, Icairns, Id447, Ideyal, Iridescence, IstvanWolf, J.delanoy, JaGa, Jac64, Jacj, JackBanks, Jaraalbe, JimVC3, Jitse Niesen, JohnOwens, Jono4174, Jossi, Jwy, Karl-Henner, Kelovy, Kent Wang, Kralizec!, Ktsquare, Kumorifox, Kurykh, Kwamikagami, Lewis R, Lisatwo, Marc Venot, MartinHarper, Materialscientist, Mav, Maximaximax, Megan1967, Michael Devore, Milkbreath, Milo99, Minesweeper, Mononomic, Montrealais, Mr3641, Mus Musculus, Mxn, Nabokov, NawlinWiki, Nergaal, Night Gyr, Nishkid64, Nk, Noclevername, Oliviosu, Oxymoron83, Paraballo, Paul Drye, Paxsimius, Physchim62, Pi zero, Plexust, Polymerbringer, Poolkris, RTC, Rebroad, Remember, Reyk, Riana, Rjwilmsi, Roadrunner, Roberta F., Romanm, S2000magician, SDC, Saperaud, Savie Kumara, Schneelocke, Scott Adler, Sengkang, Sfuerst, ShaunMacPherson, Sl, Smack, Smokefoot, Splarka, Squash, Staffwaterboy, Stan Shebs, Stifynsemons, Stone, Svante, TEDLEVITT, Tagishsimon, Tellyaddict, Tiwonk, Tjfulopp, Tom the Goober, TomaszHolband, Tristanreid, Until It Sleeps, Uzecon, V1adis1av, VASANTH S.N., Vsmith, Warut, Wireless friend, Wjmallard, Yekrats, Yyy, 220 anonymous edits
Image Sources, Licenses and Contributors image:Ta-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Ta-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Conscious, Paddy, Saperaud Image: Tantal 1.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Tantal_1.jpg License: Public Domain Contributors: Original uploader was Tomihahndorf at de.wikipedia Image:Tantalite.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Tantalite.jpg License: unknown Contributors: Tillman, 1 anonymous edits Image:Tantal-Perle-Wiki-07-02-25-P1040364b.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Tantal-Perle-Wiki-07-02-25-P1040364b.jpg License: unknown Contributors: Elcap Jens Both
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11
Tungsten
1
Tungsten 74
tantalum ← Tungsten → rhenium
Mo ↑
W ↓
Sg Periodic Table - Extended Periodic Table
General Name, Symbol, Number
Tungsten, W, 74
Element category
transition metals
Group, Period, Block
6, 6, d
Appearance
grayish white, lustrous
Standard atomic weight
183.84 (1) g·mol
Electron configuration
[Xe] 4f
−1
14
Electrons per shell
4
2[1]
5d 6s
2, 8, 18, 32, 12, 2 Physical properties
Phase
solid
Density (near r.t.)
19.25 g·cm−3
Liquid density at m.p.
17.6 g·cm−3
Melting point
3695 K (3422 °C, 6192 °F)
Boiling point
5828 K (5555 °C, 10031 °F)
Critical point
13892 K, {{{mpa}}} MPa
Heat of fusion
52.31 kJ·mol−1
Heat of vaporization
806.7 kJ·mol−1
Specific heat capacity
(25 °C) 24.27 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
3477
3773
4137
4579
5127
5823
Tungsten
2 Atomic properties
Crystal structure
body centered cubic
Oxidation states
6, 5, 4, 3, 2, 1, 0, −1, -2 (mildly acidic oxide)
Electronegativity
2.36 (Pauling scale)
Ionization energies
1st: 770 kJ/mol 2nd: 1700 kJ/mol
Atomic radius
139 pm
Covalent radius
162±7 pm Miscellaneous [2]
Magnetic ordering
paramagnetic
Electrical resistivity
(20 °C) 52.8 n Ω·m
Thermal conductivity
(300 K) 173 W·m
Thermal expansion
(25 °C) 4.5 µm·m
Speed of sound (thin rod)
(r.t.) (annealed) −1 4290 m·s
Young's modulus
411 GPa
Shear modulus
161 GPa
Bulk modulus
310 GPa
−1
−1
·K
−1
Poisson ratio
0.28
Mohs hardness
7.5
Vickers hardness
3430 MPa
Brinell hardness
2570 MPa
CAS registry number
−1
·K
7440-33-7 Most-stable isotopes
Main article: Isotopes of tungsten iso 180
W
181
W
NA 0.12% syn
half-life
DM α
2.516
176
121.2 d
ε
0.188
181
0.433
185
26.50%
W is stable with 108 neutrons
183
14.31%
W is stable with 109 neutrons
184
30.64%
W is stable with 110 neutrons
W W
185
W
186
W
syn 28.43%
DP
1.8×1018 y
182
W
DE (MeV)
75.1 d
β−
W is stable with 112 neutrons
References
Hf Ta
Re
Tungsten Tungsten (pronounced /ˈtʌŋstən/), also known as wolfram (/ˈwʊlfrəm/), is a chemical element with the chemical symbol W and atomic number 74. A steel-gray metal, tungsten is found in several ores, including wolframite and scheelite. It is remarkable for its robust physical properties, especially the fact that it has the highest melting point of all the non-alloyed metals and the second highest of all the elements after carbon.[3] Tungsten is often brittle[4] and hard to work in its raw state; however, if pure, it can be cut with a hacksaw.[5] The pure form is used mainly in electrical applications, but its many compounds and alloys are used in many applications, most notably in incandescent light bulb filaments, X-ray tubes (as both the filament and target), and superalloys. Tungsten is also the only metal from the third transition series that is known to occur in biomolecules, and is the heaviest element known to be used by living organisms.[6] [7]
History In 1781, Carl Wilhelm Scheele discovered that a new acid, tungstic acid, could be made from scheelite (at the time named tungstenite). Scheele and Torbern Bergman suggested that it might be possible to obtain a new metal by reducing this acid.[8] In 1783, José and Fausto Elhuyar found an acid made from wolframite that was identical to tungstic acid. Later that year, in Spain, the brothers succeeded in isolating tungsten by reduction of this acid with charcoal, and they are credited with the discovery of the element.[9] [10] In World War II, tungsten played a significant role in background political dealings. Portugal, as the main European source of the element, was put under pressure from both sides, because of its deposits of wolframite ore. Tungsten's resistance to high temperatures and its strength in alloys made it an important raw material for the weaponry industry.[11]
Etymology The name "tungsten" (from the Swedish, Norwegian and Danish tung sten, meaning "heavy stone") is used in English, French, Italian and many other languages as the name of the element. Tungsten was the old Swedish name for the mineral scheelite. The other name "wolfram" (or "volfram"), used for example in German, Spanish, Russian and in both Swedish and Danish, is derived from the mineral wolframite, and this is also the origin of its chemical symbol, W.[5] The name "wolframite" is derived from German "wolf rahm" ("wolf soot" or "wolf cream"), the name given to tungsten by Johan Gottschalk Wallerius in 1747. This, in turn, derives from "Lupi spuma", the name Georg Agricola used for the element in 1546, which translates into English as "wolf's froth" or "cream" (the etymology is not entirely certain), and is a reference to the large amounts of tin consumed by the mineral during its extraction.[12]
3
Tungsten
4
Characteristics Physical
Pure tungsten
In its raw form, tungsten is a steel-gray metal that is often brittle and hard to work. But, if pure, it can be worked easily.[5] It is worked by forging, drawing, extruding, or sintering. Of all metals in pure form, tungsten has the highest melting point (3,422 °C, 6,192 °F), lowest vapor pressure and (at temperatures above 1,650 °C, 3,002 °F) the highest tensile strength.[13] Tungsten has the lowest coefficient of thermal expansion of any pure metal. Alloying small quantities of tungsten with steel greatly increases its toughness.[3]
Isotopes Naturally occurring tungsten consists of five isotopes whose half-lives are so long that they can be considered stable. Theoretically, all five can decay into isotopes of element 72 (hafnium) by alpha emission, but only 180W has been observed [14] to do so with a half-life of (1.8 ± 0.2)·1018 yr; on average, this yields about two alpha decays of 180W in one gram of natural tungsten per year.[15] The other naturally occurring isotopes have not been observed to decay, constraining their half-lives to be:[15] 182
W, T1/2 > 8.3·1018 years
183
W, T1/2 > 29·1018 years
184
W, T1/2 > 13·1018 years
186
W, T1/2 > 27·1018 years
Another 30 artificial radioisotopes of tungsten have been characterized, the most stable of which are 181W with a half-life of 121.2 days, 185W with a half-life of 75.1 days, 188W with a half-life of 69.4 days, 178W with a half-life of 21.6 days, and 187W with a half-life of 23.72 h.[15] All of the remaining radioactive isotopes have half-lives of less than 3 hours, and most of these have half-lives below 8 minutes.[15] Tungsten also has 4 meta states, the most stable being 179mW (T½ 6.4 minutes).
Chemical
References [1] " Why does Tungsten not 'Kick' up an electron from the s sublevel ? (http:/ / www. madsci. org/ posts/ archives/ 2000-02/ 951518136. Ch. r. html)". . Retrieved 2008-06-15. [2] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [3] Daintith, John (2005). Facts on File Dictionary of Chemistry, 4th ed.. New York: Checkmark Books. [4] Lassner, Erik; Schubert, Wolf-Dieter (1999). " low temperature brittleness (http:/ / books. google. de/ books?id=foLRISkt9gcC& pg=PA20)". Tungsten: properties, chemistry, technology of the element, alloys, and chemical compounds. Springer. p. 256. ISBN 9780306450532. . [5] Stwertka, Albert (2002). A Guide to the elements, 2nd ed.. New York: Oxford University Press.
Tungsten
5
[6] J McMaster and John H Enemark (1998). "The active sites of molybdenum- and tungsten-containing enzymes". Current Opinion in Chemical Biology 2 (2): 201–207. doi: 10.1016/S1367-5931(98)80061-6 (http:/ / dx. doi. org/ 10. 1016/ S1367-5931(98)80061-6). [7] Russ Hille (2002). "Molybdenum and tungsten in biology". Trends in Biochemical Sciences 27 (7): 360–367. doi: 10.1016/S0968-0004(02)02107-2 (http:/ / dx. doi. org/ 10. 1016/ S0968-0004(02)02107-2). [8] Saunders, Nigel (February 2004). Tungsten and the Elements of Groups 3 to 7 (The Periodic Table). Chicago, Illinois: Heinemann Library. ISBN 1403435189. [9] " ITIA Newsletter (http:/ / www. itia. info/ FileLib/ ITIA_Newsletter_June05. pdf)" (PDF). International Tungsten Industry Association. June 2005. . Retrieved 2008-06-18. [10] " ITIA Newsletter (http:/ / www. itia. info/ FileLib/ ITIA_Newsletter_December05. pdf)" (PDF). International Tungsten Industry Association. December 2005. . Retrieved 2008-06-18. [11] Stevens, Donald G. (1999). " World War II Economic Warfare: The United States, Britain, and Portuguese Wolfram (http:/ / www. questia. com/ googleScholar. qst;jsessionid=LY1PyzmCc1D256Gvh5wpbhxKyTyvcm2FHpMwpcs2wW2XyytCh4pW!956463030?docId=5001286099)". The Historian ( Questia (http:/ / www. questia. com)). . [12] Peter van der Krogt. " Wolframium Wolfram Tungsten (http:/ / elements. vanderkrogt. net/ elem/ w. html)". Elementymology & Elements Multidict. . Retrieved 2008-05-09. [13] C. R. Hammond (2004). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304857. [14] C. Cozzini et al. (2004). " Detection of the natural α decay of tungsten (http:/ / arxiv. org/ abs/ nucl-ex/ 0408006)". Phys. Rev. C 70: 064606. doi: 10.1103/PhysRevC.70.064606 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 70. 064606). . [15] Alejandro Sonzogni. " Interactive Chart of Nuclides (http:/ / www. nndc. bnl. gov/ chart/ )". Brookhaven National Laboratory. . Retrieved 2008-06-06. [16] Emsley, John E. (1991). The elements, 2nd ed.. New York: Oxford University Press. [17] Smith, Bradley J. (2000). " Quantitative Determination of Sodium Metatungstate Speciation by 183W N.M.R. Spectroscopy (http:/ / www. publish. csiro. au/ paper/ CH00140. htm)". Australian Journal of Chemistry (CSIRO) 53 (12). . Retrieved 2008-06-17. [18] Lassner, Erik (1999). Tungsten: Properties, Chemistry, Technology of the Element, Alloys and Chemical Compounds (http:/ / books. google. com/ books?id=foLRISkt9gcC& pg=PA409& lpg=PA409& dq=tungsten+ nutrient+ organisms& source=web& ots=-rtHF9sWBY& sig=CoCD7Wp0HS-QRzQEoiPCisLaP04& hl=en& sa=X& oi=book_result& resnum=1& ct=result). Springer. pp. 409–411. ISBN 0306450534. . [19] Mullen, Frank X. (April 27, 2006). " Mouse Study Findings key in Fallon Cancer Cases, Scientists Say (http:/ / www. familiesagainstcancer. org/ ?id=344)". Reno Gazette-Journal. . Retrieved 2008-06-17. [20] Shedd, Kim B. (2000). " Tungsten (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ tungsten/ 680400. pdf)" (PDF). United States Geological Survey. . Retrieved 2008-06-18. [21] Schey, John A. (1987). Introduction to Manufacturing Processes, 2nd ed.. McGraw-Hill, Inc. [22] " US Patent 6428904 - X-ray target (http:/ / www. patentstorm. us/ patents/ 6428904/ description. html)". PatentStorm (http:/ / www. patentstorm. us/ ). August 6, 2002. . Retrieved 2008-06-18. [23] " Tungsten Applications - Steel (http:/ / www. azom. com/ details. asp?ArticleID=1264)". azom.com (http:/ / www. azom. com/ ). 2000-2008. . Retrieved 2008-06-18. [24] Gray, Theo (March 14, 2008). " How to Make Convincing Fake-Gold Bars (http:/ / www. popsci. com/ diy/ article/ 2008-03/ how-make-convincing-fake-gold-bars)". Popular Science. . Retrieved 2008-06-18.
Article Sources and Contributors
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6
Rhenium
1
Rhenium tungsten ← rhenium → osmium Tc ↑ Re ↓ Bh Periodic Table Extended Periodic Table General Name, symbol, number
rhenium, Re, 75
Element category
transition metals
Group, Period, Block
7, 6, d
Appearance
grayish white
Standard atomic weight
−1
186.207(1) g·mol
Electron configuration
14
[Xe] 4f
Electrons per shell
5
2
5d 6s
2, 8, 18, 32, 13, 2 Physical properties
Phase
solid Density (near r.t.)
21.02 g·cm−3
Liquid density at m.p.
18.9 g·cm−3
Melting point
3459 K (3186 °C, 5767 °F) Boiling point
5869 K (5596 °C, 10105 °F)
Heat of fusion
60.43 kJ·mol−1
Heat of vaporization
704 kJ·mol−1
Specific heat capacity
(25 °C) 25.48 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
3303
3614
4009
4500
5127
5954
Atomic properties
Rhenium
2 Crystal structure
hexagonal
Oxidation states
7, 6, 5, 4, 3, 2, 1, 0, -1 (mildly acidic oxide)
Electronegativity
1.9 (Pauling scale)
Ionization energies (more)
1st: 760 kJ·mol−1 2nd: 1260 kJ·mol−1 3rd: 2510 kJ·mol−1
Atomic radius
137 pm
Covalent radius
151±7 pm Miscellaneous [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(20 °C) 193 n Ω·m
Thermal conductivity
(300 K) 48.0 W·m
Thermal expansion
(25 °C) 6.2 µm·m
Speed of sound (thin rod)
(20 °C) 4700 m/s
Young's modulus
463 GPa
Shear modulus
178 GPa Bulk modulus
−1
−1
−1
−1
·K
·K
370 GPa
Poisson ratio
0.30
Mohs hardness
7.0
Vickers hardness
2450 MPa
Brinell hardness
1320 MPa
CAS registry number
7440-15-5 Most-stable isotopes
Main article: Isotopes of rhenium iso
NA
half-life
185
37.4%
185
187
62.6%
4.35×1010 y
Re Re
DM
DE (MeV)
DP
Re is stable with 110 neutron α (not observed)
1.653
183
β-
0.0026
187
Ta Os
References
Rhenium (pronounced /ˈriːniəm/) is a chemical element with the symbol Re and atomic number 75. It is a silvery-white, heavy, third-row transition metal in group 7 of the periodic table. With an average concentration of 1 part per billion (ppb), rhenium is one of the rarest elements in the Earth's crust. Rhenium resembles manganese chemically and is obtained as a by-product of molybdenum and copper refinement. Rhenium shows in its compounds a wide variety of oxidation states ranging from −1 to +7.
Rhenium Minor amounts of rhenium are added into tungsten alloys and some rhenium compounds are used as catalysts in the chemical industry. Nickel-based superalloys for the use in jet engines contain up to 6% of rhenium, making it the largest use for rhenium. Because of the low availability and the demand for jet engines, rhenium is among the most expensive metals on Earth, whose price at times exceeds US$12,000 per kilogram. Rhenium, discovered in 1925, was the last naturally occurring stable element to be discovered. Francium was the last identified naturally-occurring element, but it is unstable. Rhenium was named after the river Rhine.
History Rhenium (Latin Rhenus meaning Rhine)[2] was the next-to-last naturally occurring element to be discovered and the last element to be discovered having a stable isotope.[3] The existence of a yet undiscovered element at this position in the periodic table had been predicted by Henry Moseley in 1914.[4] It is generally considered to have been discovered by Walter Noddack, Ida Tacke, and Otto Berg in Germany. In 1925 they reported that they detected the element in platinum ore and in the mineral columbite. They also found rhenium in gadolinite and molybdenite.[5] In 1928 they were able to extract 1 g of the element by processing 660 kg of molybdenite.[6] The process was so complicated and expensive that production was discontinued until early 1950 when tungsten-rhenium and molybdenum-rhenium alloys were prepared. These alloys found important applications in industry that resulted in a great demand for the rhenium produced from the molybdenite fraction of porphyry copper ores. In 1908, Japanese chemist Masataka Ogawa announced that he discovered the 43rd element and named it nipponium (Np) after Japan (which is Nippon in Japanese). However, later analysis indicated the presence of rhenium (element 75), not element 43.[7] The symbol Np was later used for the element neptunium.
Characteristics Rhenium is a silvery-white metal with one of the highest melting points of all elements, exceeded by only tungsten and carbon. It is also one of the densest, exceeded only by platinum, iridium and osmium. Its usual commercial form is a powder, but this element can be consolidated by pressing and sintering in a vacuum or hydrogen atmosphere. This procedure yields a compact having the density above 90% of the density of the metal. When annealed this metal is very ductile and can be bent, coiled, or rolled.[8] Rhenium-molybdenum alloys are superconductive at 10 K; tungsten-rhenium alloys are also superconductive[9] around 4-8 K, depending on the alloy. Rhenium metal superconducts at 2.4 K.[10] [11]
Isotopes Naturally occurring rhenium is 37.4% 185Re, which is stable, and 62.6% 187Re, which is unstable but has a very long half-life (~1010 years); that lifetime is affected by the charge state of rhenium atom[12] [13] . The beta decay of 187Re is used for rhenium-osmium dating of ores. The available energy for this beta decay (2.6 keV) is one of the lowest known among all radionuclides. There are twenty-six other recognized radioactive isotopes of rhenium.[14]
3
Rhenium
Compounds Rhenium has the widest range of oxidation states of any known element: −1, 0, +1, +2, +3, +4, +5, +6 and +7.[15] The oxidation states +7, +6, +4, and +2 are the most common.[15] The most common rhenium compounds are the oxides and the halides exhibiting a broad oxidation number spectrum: Re2O7, ReO3, Re2O5, ReO2, and Re2O3 are the known oxides, and ReF7, ReCl6, ReCl5, ReCl4 and ReCl3 are a few of the known halogen derivatives.[16] Known sulfides are ReS2 and Re2S7.[16] Rhenium is most available commercially as the sodium and ammonium perrhenates. It is also readily available as dirhenium decacarbonyl; these three compounds are common entry points to rhenium chemistry. Various perrhenate salts may be easily converted to tetrathioperrhenate by the action of ammonium hydrosulfide.[17] It is possible to reduce the dirhenium decacarbonyl Re2(CO)10 by reacting it with sodium amalgam to Na[Re(CO)5] with rhenium in the formal oxidation state −1.[18] Dirhenium decacarbonyl may be oxidatively cleaved with bromine to give bromopentacarbonylrhenium(I),[19] then reduced with zinc and acetic acid to pentacarbonylhydridorhenium:[20] Re2(CO)10 + Br2 → Re(CO)5Br Re(CO)5Br + Zn + HOAc → Re(CO)5H + ZnBr(OAc) Bromopentacarbonylrhenium(I) may be decarbonylated to give the rhenium tricarbonyl fragment either by refluxing in water:[21] Re(CO)5Br + 3 H2O → [Re(CO)3(H2O)3]Br + 2 CO or by reacting with tetraethylammonium bromide:[22] Re(CO)5Br + 2 (NEt4Br → [NEt4]2[Re(CO)3Br3] Rhenium diboride (ReB2) is a hard compound having the hardness similar to that of tungsten carbide, silicon carbide, titanium diboride or zirconium diboride.[23] Rhenium was originally thought to form the rhenide anion, Re−, in which it has the −1 oxidation state. This was based on the product of the reduction of perrhenate salts, such as the reduction of potassium perrhenate (KReO4) by potassium metal.[24] "Potassium rhenide" was shown to exist as a tetrahydrated complex, with the postulated chemical formula KRe·4H2O.[25] This compound exhibits strongly reducing properties, and slowly yields hydrogen gas when dissolved in water. The lithium and thallous salts were also reported. Later research, however, indicates that the "rhenide" ion is actually a hydridorhenate complex. "Potassium rhenide" was shown to be in fact the nonahydridorhenate, K2ReH9, containing the ReH2−9 anion in which the oxidation state of rhenium is actually +7.[26] [27] Other methods of reduction of perrhenate salts yield compounds containing other hydridocomplexes, including ReH3(OH)3(H2O)−.[28]
4
Rhenium
Occurrence Rhenium is one of the rarest elements in Earth's crust with a average concentration of 1 ppb;[16] other sources quote the number of 0.5 ppb making it the 77th most abundant element in Earth's crust.[29] Rhenium is probably not found free in nature (its possible natural occurrence is uncertain), but occurs in amounts up to 0.2%[16] in the mineral molybdenite, the major commercial source, although single molybdenite samples with up to 1.88% have been found.[30] Chile Molybdenite has the world's largest rhenium reserves, part of the copper ore deposits, and was the leading producer as of [31] 2005. It was only recently that the first rhenium mineral was found and described (in 1994), a rhenium sulfide mineral (ReS2) condensing from a fumarole on Russia's Kudriavy volcano, in the Kurile Islands.[32] Named rheniite, this rare mineral commands high prices among collectors,[33] but is not an economically viable source of the element.
Production Commercial rhenium is extracted from molybdenum roaster-flue gas obtained from copper-sulfide ores. Some molybdenum ores contain 0.001% to 0.2% rhenium.[16] [30] Rhenium(VII) oxide and perrhenic acid readily dissolve in water; they are leached from flue dusts and gasses and extracted by precipitating with potassium or ammonium chloride as the perrhenate salts, and purified by recrystallization.[34] Total world Ammonium perrhenate production is between 40 and 50 tons/year; the main producers are in Chile, the United States, and Kazakhstan.[35] Recycling of used Pt-Re catalyst and special alloys allow the recovery of another 10 tons per year. Prices for the metal rose rapidly in early 2008, from $1000–$2000 per kg in 2003-2006 to over $10,000 in February 2008.[36] [37] The metal form is prepared by reducing ammonium perrhenate with hydrogen at high temperatures:[34] 2 NH4ReO4 + 7 H2 → 2 Re + 8 H2O + 2 NH3
5
Rhenium
6
Applications Rhenium is added to high-temperature superalloys that are used to make jet engine parts, making 70% of the worldwide rhenium production.[38] Another major application is in platinum-rhenium catalysts, which are primarily used in making lead-free, high-octane gasoline.[35] [39]
Alloys
The F-15 engine uses rhenium containing second-generation superalloys
The nickel-based superalloys have improved creep strength with the addition of rhenium. The alloys normally contain 3% or 6% of rhenium.[40] The second generation alloys contain 3%; these alloys were used in the engines of the F-16 and F-15, while the new single-crystal third-generation alloys contain 6% of rhenium; they are used in the F-22 and F-35 engines.[39] [41] For 2006 the consumption is given as 28% for General Electric, 28% Rolls-Royce plc and 12% Pratt & Whitney, all for superalloys, while the use for catalysts only accounts for 14% and the remaining applications use 18%.[38] In 2006, 77% of the rhenium consumption in the United States was in alloys.[39] Rhenium improves the properties of tungsten and is therefore the most important alloying material for tungsten. Tungsten-rhenium alloys are more ductile at low temperature making them easier to machine, while the high-temperature stability is also improved. The effect increases with the rhenium concentration, and therefore tungsten alloys are produced with up to 27% of Re, which is the solubility limit.[42] One application for the tungsten-rhenium alloys is x-ray sources. The high melting point of both compounds, together with the high atomic mass, makes them stable against the prolonged electron impact.[43] Rhenium tungsten alloys are also applied as thermocouples to measure temperatures up to 2200 °C.[44] The high temperature stability, low vapor pressure, good wear resistance and ability to withstand arc corrosion of rhenium are useful in self-cleaning electrical contacts. In particular, the discharge occurring during the switching oxidizes the contacts. However, rhenium oxide Re2O7 has poor stability (sublimates at ~360 °C) and therefore is removed during the discharge.[38] Rhenium has a high melting point and a low vapor pressure similar to tantalum and tungsten, however, rhenium forms no volatile oxides. Therefore, rhenium filaments exhibit a higher stability if the filament is operated not in vacuum, but in oxygen-containing atmosphere.[45] Those filaments are widely used in mass spectrographs, in ion gauges.[46] and in photoflash lamps in photography.[47]
Catalysts Rhenium in the form of rhenium-platinum alloy is used as catalyst for catalytic reforming, which is a chemical process to convert petroleum refinery naphthas with low octane ratings into high-octane liquid products. Worldwide, 30% of catalysts used for this process contain rhenium.[48] The olefin metathesis is the other reaction for which rhenium is used as catalyst. Normally Re2O7 on alumina is used for this process.[49] Rhenium catalysts are very resistant to chemical poisoning from nitrogen, sulfur and phosphorus, and so are used in
Rhenium certain kinds of hydrogenation reactions.[8]
7 [50] [51]
Other uses 188
Rh and 186Rh isotopes are radioactive and are used for treatment of liver cancer. They both have similar penetration depth in tissue (5 mm for 186Rh and 11 mm for 188Rh), but 186 Re has advantage of longer lifetime (90 hours vs. 17 hours).[52] [53] Related by periodic trends, rhenium has a similar chemistry with technetium; work done to label rhenium onto target compounds can often be translated to technetium. This is useful for radiopharmacy, where it is difficult to work with technetium - especially the 99m isotope used in medicine - due to its expense and short half-life.[52] [54]
Precaution Very little is known about the toxicity of rhenium and its compounds because they are used in very small amounts. Soluble salts, such as the rhenium halides or perrhenates, could be hazardous due to elements other than rhenium or due to rhenium itself.[55] Only a few compounds of rhenium have been tested for their toxicity; two examples are potassium perrhenate and rhenium trichloride, which were injected as a solution into rats. The perrhenate had an LD50 value of 2800 mg/kg after seven days and the rhenium trichloride showed LD50 of 280 mg/kg.[56]
External links • WebElements.com - Rhenium [57] • pure Rhenium >99,99% picture in the element collection from Heinrich Pniok
[58]
pnb:مینیہر
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] Tilgner, Hans Georg (2000) (in German). Forschen Suche und Sucht (http:/ / books. google. com/ books?id=UWBWnMOGtMQC). Books on Demand. ISBN 9783898112727. . [3] " Rhenium: Statistics and Information (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ rhenium/ )". Minerals Information. United States Geological Survey. 2008. . Retrieved 2008-02-03. [4] Moseley, Henry (1914). " High Frequency Spectra of the Elements, Part II (http:/ / www. chemistry. co. nz/ henry_moseley_article. htm)". Philosophical Magazine. . [5] Noddack, W.; Tacke, I.; Berg, O. (1925). "Die Ekamangane". Naturwissenschaften 13 (26): 567–574. doi: 10.1007/BF01558746 (http:/ / dx. doi. org/ 10. 1007/ BF01558746). [6] Noddack, W.; Noddack , I. (1929). "Die Herstellung von einem Gram Rhenium" (in German). Zeitschrift für anorganische und allgemeine Chemie 183 (1): 353–375. doi: 10.1002/zaac.19291830126 (http:/ / dx. doi. org/ 10. 1002/ zaac. 19291830126). [7] Yoshihara, H. K. (2004). "Discovery of a new element ‘nipponiumʼ: re-evaluation of pioneering works of Masataka Ogawa and his son Eijiro Ogawa". Spectrochimica Acta Part B Atomic Spectroscopy 59: 1305–1310. doi: 10.1016/j.sab.2003.12.027 (http:/ / dx. doi. org/ 10. 1016/ j. sab. 2003. 12. 027). [8] Hammond, C. R. (2004). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304857. [9] Neshpor, V. S.; Novikov, V. I.; Noskin, V. A.; Shalyt, S. S. (1968). "Superconductivity of Some Alloys of the Tungsten-rhenium-carbon System". Soviet Physics JETP 27: 13. Bibcode: 1968JETP...27...13N (http:/ / adsabs. harvard. edu/ abs/ 1968JETP. . . 27. . . 13N). [10] Daunt, J. G.; Smith, T. S. (1952). "Superconductivity of Rhenium". Physical Review 88 (2): 309–311. doi: 10.1103/PhysRev.88.309 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 88. 309).
Rhenium [11] Daunt, J. G.; Lerner, E.. " The Properties of Superconducting Mo-Re Alloys (http:/ / stinet. dtic. mil/ oai/ oai?verb=getRecord& metadataPrefix=html& identifier=AD0622881)". Defense Technical Information Center. . [12] Johnson, Bill (1993). " How to Change Nuclear Decay Rates (http:/ / math. ucr. edu/ home/ baez/ physics/ ParticleAndNuclear/ decay_rates. html)". . Retrieved 2009-02-21. [13] Bosch (1996). "Observation of bound-state β– decay of fully ionized 187Re:187Re-187Os Cosmochronometry". Physical Review Letters 77 (26): 5190–5193. doi: 10.1103/PhysRevLett.77.5190 (http:/ / dx. doi. org/ 10. 1103/ PhysRevLett. 77. 5190). [14] Georges, Audi (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3–128. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [15] Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils; (1985). "Rhenium" (in German). Lehrbuch der Anorganischen Chemie (91–100 ed.). Walter de Gruyter. pp. 1118–1123. ISBN 3110075113. [16] Woolf, A. A. (1961). "An outline of rhenium chemistry". Quarterly Review of the Chemical Society 15: 372–391. doi: 10.1039/QR9611500372 (http:/ / dx. doi. org/ 10. 1039/ QR9611500372). [17] Goodman, J. T.; Rauchfuss, T. B. (2002). "Tetraethylammonium-tetrathioperrhenate [Et4N][ReS4]". Inorganic Syntheses 33: 107–110. [18] Breimair, Josef (1990). "Nucleophile Addition von Carbonylmetallaten an kationische Alkin-Komplexe [CpL2M(η2-RC≡CR)]+ (M = Ru, Fe): μ-η1:η1-Alkin-verbrückte Komplexe". Chemische Berichte 123: 7. doi: 10.1002/cber.19901230103 (http:/ / dx. doi. org/ 10. 1002/ cber. 19901230103). [19] Schmidt, Steven P.; Trogler, William C.; Basolo, Fred (1990). "Pentacarbonylrhenium Halides". Inorganic Syntheses 28: 154–159. doi: 10.1002/9780470132593.ch42 (http:/ / dx. doi. org/ 10. 1002/ 9780470132593. ch42). [20] Michael A. Urbancic, John R. Shapley (1990). "Pentacarbonylhydridorhenium". Inorganic Syntheses 28: 165–168. doi: 10.1002/9780470132593.ch43 (http:/ / dx. doi. org/ 10. 1002/ 9780470132593. ch43). [21] Lazarova, N.; James, S.; Babich, J.; Zubieta, J. (2004). "A convenient synthesis, chemical characterization and reactivity of [Re(CO)3(H2O)3]Br: the crystal and molecular structure of [Re(CO)3(CH3CN)2Br]". Inorganic Chemistry Communications 7 (9): 1023–1026. doi: 10.1016/j.inoche.2004.07.006 (http:/ / dx. doi. org/ 10. 1016/ j. inoche. 2004. 07. 006). [22] Alberto, R.; Egli, A.; Abram, U.; Hegetschweiler, K.; Gramlich V.; Schubiger, P. A. (1994). "Synthesis and reactivity of [NEt4]2[ReBr3(CO)3]. Formation and structural characterization of the clusters [NEt4][Re3(µ3-OH)(µ-OH)3(CO)9] and [NEt4][Re2(µ-OH)3(CO)6] by alkaline titration". J. Chem. Soc., Dalton Trans.: 2815–2820. doi: 10.1039/DT9940002815 (http:/ / dx. doi. org/ 10. 1039/ DT9940002815). [23] Qin, Jiaqian; He, Duanwei; Wang, Jianghua; Fang, Leiming; Lei, Li; Li, Yongjun; Hu, Juan; Kou, Zili; Bi, Yan (2008). "Is Rhenium Diboride a Superhard Material?". Advanced Materials 20: 4780–4783. doi: 10.1002/adma.200801471 (http:/ / dx. doi. org/ 10. 1002/ adma. 200801471). [24] Cobble, J. W. (June 1957). "On the Structure of the Rhenide Ion". The Journal of Physical Chemistry 61 (6): 727–729. doi: 10.1021/j150552a005 (http:/ / dx. doi. org/ 10. 1021/ j150552a005). edit (http:/ / en. wikipedia. org/ wiki/ Template:cite_doi/ 10. 1021. 2fj150552a005) [25] Bravo, Justo B.; Ernest Griswold; Jacob Kleinberg (January 1954). "The Preparation of a Solid Rhenide". The Journal of Physical Chemistry 58 (1): 18–21. doi: 10.1021/j150511a004 (http:/ / dx. doi. org/ 10. 1021/ j150511a004). edit (http:/ / en. wikipedia. org/ wiki/ Template:cite_doi/ 10. 1021. 2fj150511a004) [26] Floss, J. G.; Grosse, A. V. (1960). "Alkali and alkaline earth rhenohydrides". Journal of Inorganic and Nuclear Chemistry 16: 36–43. doi: 10.1016/0022-1902(60)80083-8 (http:/ / dx. doi. org/ 10. 1016/ 0022-1902(60)80083-8). edit (http:/ / en. wikipedia. org/ wiki/ Template:cite_doi/ 10. 1016. 2f0022-1902. 2860. 2980083-8) [27] Kenneth Malcolm Mackay; Rosemary Ann Mackay; W. Henderson (2002). Rosemary Ann Mackay. ed. Introduction to modern inorganic chemistry (6th ed.). CRC Press. pp. 368–369. ISBN 0748764208. [28] M. L. H. Green; D. J. Jones (1965). H.J. Emeleus, A.G. Sharpe. ed. Advances in inorganic chemistry and radiochemistry. Academic Press. pp. 169–172. ISBN 0120236079. [29] Emsley, John (2001). " Rhenium (http:/ / books. google. com/ books?id=j-Xu07p3cKwC)". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. pp. 358–360. ISBN 0-19-850340-7. . [30] Rouschias, George (1974). "Recent advances in the chemistry of rhenium". Chemical Reviews 74: 531. doi: 10.1021/cr60291a002 (http:/ / dx. doi. org/ 10. 1021/ cr60291a002). [31] Anderson, Steve T. " 2005 Minerals Yearbook: Chile (http:/ / minerals. usgs. gov/ minerals/ pubs/ country/ 2005/ cimyb05. pdf)" (PDF). United States Geological Survey. . Retrieved 2008-10-26. [32] Korzhinsky, M.A.; Tkachenko, S. I.; Shmulovich, K. I.; Taran Y. A.; Steinberg, G. S. (2004-05-05). "Discovery of a pure rhenium mineral at Kudriavy volcano". Nature 369: 51–52. doi: 10.1038/369051a0 (http:/ / dx. doi. org/ 10. 1038/ 369051a0).
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Rhenium [33] " The Mineral Rheniite (http:/ / www. galleries. com/ minerals/ sulfides/ rheniite/ rheniite. htm)". Amethyst Galleries,Inc.. . [34] Patnaik, Pradyot (2003). Handbook of Inorganic Chemicals. McGraw-Hill. pp. 790. ISBN 0070494398. OCLC 47726843 (http:/ / worldcat. org/ oclc/ 47726843). [35] Magyar, Michael J. (January 2008). " Rhenium (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ rhenium/ mcs-2008-rheni. pdf)" (PDF). Mineral Commodity Summaries. U.S. Geological Survey. . Retrieved 2008-02-17. [36] " MinorMetal prices (http:/ / www. minormetals. com/ )". minormetals.com. . Retrieved 2008-02-17. [37] Harvey, Jan (2008-07-10). " Analysis: Super hot metal rhenium may reach "platinum prices" (http:/ / in. reuters. com/ article/ oilRpt/ idINL1037587920080710)". Reuters India. . Retrieved 2008-10-26. [38] Naumov, A. V. (2007). "Rhythms of rhenium". Russian Journal of Non-Ferrous Metals 48 (6): 418–423. doi: 10.3103/S1067821207060089 (http:/ / dx. doi. org/ 10. 3103/ S1067821207060089). [39] Magyar, Michael J.. " Mineral Yearbook: Rhenium (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ rhenium/ myb1-2006-rheni. pdf)" (PDF). United States Geological Survey. . [40] Bhadeshia, H. K. D. H.. " Nickel Based Superalloys (http:/ / www. msm. cam. ac. uk/ phase-trans/ 2003/ Superalloys/ superalloys. html)". University of Cambridge. . Retrieved 2008-10-17. [41] Cantor, B.; Grant, Patrick Assender Hazel (2001). Aerospace Materials: An Oxford-Kobe Materials Text (http:/ / books. google. de/ books?id=n09-HajhRHYC). CRC Press. pp. 82–83. ISBN 9780750307420. . [42] Lassner, Erik; Schubert, Wolf-Dieter (1999). Tungsten: properties, chemistry, technology of the element, alloys, and chemical compounds (http:/ / books. google. de/ books?id=foLRISkt9gcC& pg=PA256). Springer. p. 256. ISBN 9780306450532. . [43] Cherry, Pam; Duxbury, Angela (1998). Practical radiotherapy physics and equipment (http:/ / books. google. de/ books?id=5WIBbmmDm-gC& pg=PA55). Cambridge University Press. p. 55. ISBN 9781900151061. . [44] Asamoto, R.; Novak, P. E. (1968). " Tungsten-Rhenium Thermocouples for Use at High Temperatures (http:/ / link. aip. org/ link/ ?RSINAK/ 39/ 1233/ 1)". Review of Scientific Instruments 39: 1233. doi: 10.1063/1.1683642 (http:/ / dx. doi. org/ 10. 1063/ 1. 1683642). . [45] Blackburn, Paul E. (1966). "The Vapor Pressure of Rhenium". The Journal of Physical Chemistry 70: 311–312. doi: 10.1021/j100873a513 (http:/ / dx. doi. org/ 10. 1021/ j100873a513). [46] Earle, G. D.; Medikonduri, R.; Rajagopal, N.; Narayanan, V.; Roddy, P. A. (2005). "Tungsten-Rhenium Filament Lifetime Variability in Low Pressure Oxygen Environments". IEEE Transactions on Plasma Science 33 (5): 1736–1737. doi: 10.1109/TPS.2005.856413 (http:/ / dx. doi. org/ 10. 1109/ TPS. 2005. 856413). [47] Ede, Andrew (2006). The chemical element: a historical perspective. Greenwood Publishing Group. ISBN 9780313333040. [48] Ryashentseva, Margarita A. (1998). "Rhenium-containing catalysts in reactions of organic compounds". Russian Chemical Reviews 67: 157–177. doi: 10.1070/RC1998v067n02ABEH000390 (http:/ / dx. doi. org/ 10. 1070/ RC1998v067n02ABEH000390). [49] Mol, Johannes C. (1999). "Olefin metathesis over supported rhenium oxide catalysts". Catalysis Today 51 (2): 289–299. doi: 10.1016/S0920-5861(99)00051-6 (http:/ / dx. doi. org/ 10. 1016/ S0920-5861(99)00051-6). [50] Angelidis, T. N.; Rosopoulou, D. Tzitzios V. (1999). "Selective Rhenium Recovery from Spent Reforming Catalysts". Ind. Eng. Chem. Res. 38 (5): 1830–1836. doi: 10.1021/ie9806242 (http:/ / dx. doi. org/ 10. 1021/ ie9806242). [51] Burch, Robert (1978). " The Oxidation State of Rhenium and Its Role in Platinum-Rhenium (http:/ / www. platinummetalsreview. com/ pdf/ pmr-v22-i2-057-060. pdf)" (PDF). Platinum Metals Review 22 (2): 57–60. . [52] Dilworth, Jonathan R.; Parrott, Suzanne J. (1998). "The biomedical chemistry of technetium and rhenium". Chemical Society Reviews 27: 43–55. doi: 10.1039/a827043z (http:/ / dx. doi. org/ 10. 1039/ a827043z). [53] " The Tungsten-188 and Rhenium-188 Generator Information (http:/ / www. ornl. gov/ sci/ nuclear_science_technology/ nu_med/ 188info. htm)". Oak Ridge National Laboratory. 2005. . Retrieved 2008-02-03. [54] Colton, R.; Peacock R. D. (1962). "An outline of technetium chemistry". Quarterly Reviews Chemical Society 16: 299–315. doi: 10.1039/QR9621600299 (http:/ / dx. doi. org/ 10. 1039/ QR9621600299). [55] Emsley, J. (2003). "Rhenium". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. pp. 358–361. ISBN 0198503407. [56] Haley, Thomas J.; Cartwright, Frank D. (1968). "Pharmacology and toxicology of potassium perrhenate and rhenium trichloride". Journal of Pharmaceutical Sciences 57 (2): 321–323. doi: 10.1002/jps.2600570218 (http:/ / dx. doi. org/ 10. 1002/ jps. 2600570218). [57] http:/ / www. webelements. com/ webelements/ elements/ text/ Re/ index. html [58] http:/ / www. pse-mendelejew. de/ bilder/ re. jpg
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Article Sources and Contributors
Article Sources and Contributors Rhenium Source: http://en.wikipedia.org/w/index.php?oldid=308575004 Contributors: Achim1999, Adashiel, Ahoerstemeier, Ahruman, Alexfusco5, AlimanRuna, Andrew Kanode, Andrewa, Archimerged, Arkuat, Beetstra, Benbest, BlueEarth, Borislav Dopudja, Bryan Derksen, Bushytails, CYD, Canadian-Bacon, CanisRufus, Canley, Carnildo, Chuck Y, Conversion script, Corbon, Crystal whacker, David Latapie, DavidBailey, Deli nk, Dentren, DesmanaMoschata, Dmstann, Droog Andrey, Dwmyers, EPO, ESkog, Eastlaw, Edgar181, El C, Emperorbma, Eras-mus, Erik Zachte, Eudialytos, Fedayee, Femto, Fivemack, Gene Nygaard, Greatpatton, Grendelkhan, Hadal, Hak-kâ-ngìn, Halibutt, Haxologist, Helge Skjeveland, Hellbus, Icairns, Ideyal, Jaeger5432, Jaraalbe, Jeff3000, Jerry, Junglecat, Karlhahn, Kelovy, Killiondude, Kwamikagami, LA2, LarryMorseDCOhio, Ledelste, Lewis R, LuigiManiac, Marc Venot, Materialscientist, Matthew Yeager, Mav, Mgdurand, Minesweeper, Mr Stephen, Mr. Billion, Mr0t1633, Nergaal, Nickstuckert, Niczar, Nihiltres, Persian Poet Gal, Physchim62, PlatinumX, Plexust, Pmsyyz, Poolkris, Potatoswatter, Quadell, RTC, Reinyday, Remember, Res2216firestar, Reywas92, Rifleman 82, Rjwilmsi, Roberta F., Romanm, Rossnorman, RunOrDie, Saperaud, Sasuke Sarutobi, ScaldingHotSoup, Schneelocke, Seanisaweird, Sekom, Sengkang, Sfuerst, Shawn81, Sidonuke, SkyLined, Sl, SpookyMulder, Stebulus, Stephenb, Stone, Svante, Tagishsimon, Tetracube, TexasAndroid, TheGoblin, Thingg, TimVickers, Tophe67, V1adis1av, Vicarious, Volland, Vsmith, Walkerma, Warut, Wdanwatts, WhatamIdoing, Wimt, Yakisoba, Yekrats, Yyy, Zappa711, 169 anonymous edits
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10
Osmium
1
Osmium rhenium ← osmium → iridium Ru ↑ Os ↓ Hs Periodic Table Extended Periodic Table General Name, symbol, number
osmium, Os, 76
Element category
transition metals
Group, Period, Block
8, 6, d
Appearance
silvery, blue cast
Standard atomic weight
190.23(3) g·mol
Electron configuration
[Xe] 4f14 5d6 6s2
Electrons per shell
−1
2, 8, 18, 32, 14, 2 Physical properties
Phase
solid Density (near r.t.) Liquid density at m.p.
Melting point
22.61 g·cm−3 20 g·cm−3 3306 K (3033 °C, 5491 °F)
Boiling point
5285 K (5012 °C, 9054 °F)
Heat of fusion
57.85 kJ·mol−1
Heat of vaporization
738 kJ·mol−1
Specific heat capacity
(25 °C) 24.7 J·mol−1·K−1
Osmium
2
Vapor pressure P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
3160
3423
3751
4148
4638
5256
Atomic properties Crystal structure Oxidation states
hexagonal 8, 7, 6, 5, 4, 3, 2, 1, 0, -1, -2 (mildly acidic oxide)
Electronegativity Ionization energies
2.2 (Pauling scale) 1st: 840 kJ/mol 2nd: 1600 kJ/mol
Atomic radius Covalent radius
135 pm 144±4 pm Miscellaneous
Magnetic ordering Electrical resistivity
paramagnetic
[1]
(0 °C) 81.2 nΩ·m
Thermal conductivity
(300 K) 87.6 W·m
Thermal expansion
(25 °C) 5.1 µm·m
Speed of sound (thin rod)
(20 °C) 4940 m/s
Shear modulus
222 GPa
Poisson ratio
0.25 Bulk modulus
Mohs hardness
462 GPa 7.0
Brinell hardness CAS registry number
3920 MPa 7440-04-2 Most-stable isotopes
−1
−1
−1
−1
·K
·K
Osmium
3
Main article: Isotopes of osmium iso 184
Os
185
Os
NA 0.02%
syn
half-life >5.6×1013 y (not observed) 93.6 d
186
1.59%
2.0×1015 y
187
1.96%
187
188
13.24%
188
189
16.15%
189
190
26.26%
190
Os Os Os Os Os
191
Os
192
Os
syn 40.78%
DM
1.452
184
ε
1.013
185
α
2.822
182
−
0.314
191
− −
0.414
192
−
1.141
193
−
0.097
194
W
Re W
Os is stable with 111 neutron Os is stable with 112 neutron Os is stable with 113 neutron Os is stable with 114 neutron
15.4 d
β 12
β β
>9.8×10 y (not observed)
syn
30.11 d
β
194
syn
6y
β
Os
DP
εε
193
Os
DE (MeV)
Ir Pt
Ir Ir
References
Osmium (pronounced /ˈɒzmiəm/) is a chemical element that has the symbol Os and atomic number 76. Osmium is a hard, brittle, blue-gray or blue-black transition metal in the platinum family, and is the densest natural element. The density of osmium is , slightly greater than that of iridium, the second densest element. Osmium is found in nature as an alloy, mostly in platinum ores. Osmium is also used in alloys, with platinum, iridium and other platinum group metals. Those alloys are employed in fountain pen tips, electrical contacts and in other applications where extreme durability and hardness are needed.
Characteristics Physical Osmium is an extremely dense, blue-gray, hard but brittle metal that remains lustrous even at high temperatures. Due to its hardness, brittleness, and very high melting point (the tenth highest of all elements), solid osmium is difficult to machine, form, or work. Osmium is generally considered to be the densest known element, narrowly defeating iridium.[2] Calculations of density from the space lattice may produce the most reliable data for these elements, giving a density of for iridium versus for osmium.[3] The extraordinary density of osmium is a consequence of the lanthanide contraction.[3] Osmium possesses quite remarkable chemical and physical properties. It has the highest melting point and the lowest vapor pressure in the platinum family. Osmium has a very low compressibility. Correspondingly, its bulk modulus is extremely high, reported between and , which rivals that of diamond (). However, the hardness of osmium is low, only .[4] [5] [6]
Osmium
4
Chemical Oxidation states of osmium -2
Na2[Os(CO)4]
-1
Na2[Os4(CO)13]
0
Os3(CO)12
+1
OsI
+2
OsI2
+3
OsBr3
+4
OsO2, OsCl4
+5
OsF5
+6
OsF6
+7
OsOF5
+8
OsO4
Osmium forms compounds with the oxidation states ranging from -2 to +8. The most common oxidation states are +2, +3, +4, and +8. The +8 oxidation state is notable for being the highest attained by any chemical element, and aside from osmium, is encountered only in xenon and ruthenium.[7] The oxidation state -1 and -2 represented by the two reactive compounds Na2[Os4(CO)13] and Na2[Os(CO)4] are used in the synthesis of osmium cluster compounds.[8] [9] The most common compound exhibiting the +8 oxidation state is osmium tetroxide. This toxic compound is formed when powdered osmium is exposed to air, and is a very volatile, water-soluble, pale yellow, crystalline solid with a strong smell. Therefore, osmium powder has a characteristic smell of osmium tetroxide.[10] Osmium tetroxide forms red osmates [OsO4(OH)2]2− upon reaction with a base. With ammonia, it forms the nitrido-osmates [OsO3N]−.[11] [12] [13] Osmium tetroxide boils at 130 °C and is a powerful oxidizing agent. By contrast, osmium dioxide (OsO2) is black, non-volatile, and much less reactive and toxic. Only two osmium compounds have major applications: osmium tetroxide — for staining tissue in electron microscopy and the non-volatile osmates for organic oxidation reactions.[14] Osmium heptafluoride (OsF7) and osmium pentafluoride (OsF5) are known, but osmium trifluoride (OsF3) has not been synthesized yet. The lower oxidation states are stabilized by the larger halogens. Therefore, the trichloride, tribromide, triiodide and even osmium diiodide are known. The oxidation state +1 is only known for the osmium iodide (OsI), whereas several carbonyl complexes of osmium, such as triosmium dodecacarbonyl (Os3(CO)12), represent the oxidation state 0.[11] [12] [15] [16] In general, the lower oxidation states of osmium are stabilized by ligands that are good σ-donors (such as amines) and π-acceptors (heterocycles containing nitrogen). The higher oxidation states are stabilized by strong σ- and π-donors, such as O2− and N3−.[17]
Osmium
Isotopes Osmium has seven naturally occurring isotopes, six of which are stable: 184Os, 187Os, 188Os, 189 Os, 190Os, and (most abundant) 192Os. 186Os undergoes alpha decay with such long half-life ((2.0±1.1) × 1015 years) that for practical purposes it can be considered stable. Alpha decay is predicted for all 7 naturally occurring isotopes, but due to very long half-lives, it was observed only for 186Os. It is predicted that 184Os and 192Os can undergo double beta decay but this radioactivity has not been observed yet.[18] 187
Os is the daughter of 187Re (half-life ) and is used extensively in dating terrestrial as well as meteoric rocks (see rhenium-osmium dating). It has also been used to measure the intensity of continental weathering over geologic time and to fix minimum ages for stabilization of the mantle roots of continental cratons. This decay is a reason why rhenium-rich minerals are abnormally rich in 187Os.[19] However, the most notable application of Os in dating has been in conjunction with iridium, to analyze the layer of shocked quartz along the K-T boundary that marks the extinction of the dinosaurs 65 million years ago.[20]
History Osmium (from Greek osme meaning "smell") was discovered in 1803 by Smithson Tennant and William Hyde Wollaston in London, England.[21] The discovery of osmium is intertwined with that of platinum and the other metals of the platinum group. Platinum reached Europe as platina ("small silver"), first encountered in the late 17th century in silver mines around the Chocó Department, in Colombia.[22] The discovery that this metal was not an alloy, but a distinct new element, was published in 1748.[23] Chemists who studied platinum dissolved it in aqua regia (a mixture of hydrochloric and nitric acids) to create soluble salts. They always observed a small amount of a dark, insoluble residue.[24] Joseph Louis Proust thought that the residue was graphite.[24] Victor Collet-Descotils, Antoine François, comte de Fourcroy, and Louis Nicolas Vauquelin also observed the black residue in 1803, but did not obtain enough material for further experiments.[24] In 1803, Smithson Tennant analyzed the insoluble residue and concluded that it must contain a new metal. Vauquelin treated the powder alternately with alkali and acids[25] and obtained a volatile new oxide, which he believed to be of this new metal—which he named ptene, from the Greek word πτηνος (ptènos) for winged.[26] [27] However, Tennant, who had the advantage of a much larger amount of residue, continued his research and identified two previously undiscovered elements in the black residue, iridium and osmium.[24] [25] He obtained a yellow solution (probably of cis–[Os(OH)2O4]2−) by reactions with sodium hydroxide at red heat. After acidification he was able to distill the formed OsO4.[26] He named osmium after Greek osme meaning "a smell", because of the smell of the volatile osmium tetroxide.[28] Discovery of the new elements was documented in a letter to the Royal Society on June 21, 1804.[24] [29] Uranium and osmium were early successful catalysts in the Haber process, the nitrogen fixation reaction of nitrogen and hydrogen to produce ammonia, giving enough yield to make the process economically successful. However, in 1908 cheaper catalysts based on iron and iron oxides were introduced for the first pilot plants.[30] Nowadays, osmium is primarily obtained from the processing of platinum and nickel ores.[31]
5
Osmium
Occurrence Osmium is one of the least abundant elements in the Earth's crust with an average mass fraction of 0.05 ppb in the continental crust.[32] Osmium is found in nature as an uncombined element or in natural alloys; especially the iridium–osmium alloys, osmiridium (osmium rich), and iridiosmium (iridium rich).[25] In the nickel and copper deposits, the platinum group metals occur as sulfides (i.e. (Pt,Pd)S)), tellurides (e.g. PtBiTe), antimonides (e.g. PdSb), and arsenides (e.g. PtAs2); in all these compounds platinum is exchanged by a small amount of iridium and osmium. As with all of the platinum group metals, osmium can be found naturally in alloys with nickel or copper.[33] Within the Earth's crust, osmium, like iridium, is found at highest concentrations in three types of geologic structure: igneous deposits (crustal intrusions from below), impact craters, and deposits reworked from one of the former structures. The largest known primary reserves are in the Bushveld igneous complex in South Africa,[34] though the large copper–nickel deposits near Native platinum containing traces of Norilsk in Russia, and the Sudbury Basin in Canada are the other Platinum group metals also significant sources of osmium. Smaller reserves can be found in the United States.[34] The alluvial deposits used by pre-Columbian people in the Chocó Department, Colombia are still a source for platinum group metals. The second large alluvial deposit was found in the Ural Mountains, Russia, which is still mined.[31] [35]
Production Osmium is obtained commercially as a by-product from nickel and copper mining and processing. During electrorefining of copper and nickel, noble metals such as silver, gold and the platinum group metals, together with non-metallic elements such as selenium and tellurium settle to the bottom of the cell as anode mud, which forms the starting material for their extraction.[36] [37] In order to separate the metals, they must first be brought into solution. Several methods are available depending on the separation process and the composition of the mixture; two representative methods are fusion with sodium peroxide followed by dissolution Osmium bead in aqua regia, and dissolution in a mixture of chlorine with hydrochloric acid.[34] [38] Osmium, ruthenium, rhodium and iridium can be separated from platinum, gold and base metals by their insolubility in aqua regia, leaving a solid residue. Rhodium can be separated from the residue by treatment with molten sodium bisulphate. The insoluble residue, containing Ru, Os and Ir, is treated with sodium oxide, in which Ir is insoluble, producing water-soluble Ru and Os salts. After oxidation to the volatile oxides, RuO4 is separated from OsO4 by precipitation of (NH4)3RuCl6 with ammonium chloride. After it is dissolved, osmium is separated from the other platinum group metals by distillation or extraction with organic solvents of the volatile osmium tetroxide.[39] The first
6
Osmium
7
method is similar to the procedure used by Tennant and Wollastone. Both methods are suitable for industrial scale production. In either case, the product is reduced using hydrogen, yielding the metal as a powder or sponge that can be treated using powder metallurgy techniques.[40] Neither the producers nor the United States Geological Survey published any production amounts for osmium. Estimations of the United States consumption date published from 1971,[41] which gives a consumption in the United States of 2000 troy ounces (62 kg), would suggest that the production is still less than 1 ton per year.
Applications Because of the volatility and extreme toxicity of its oxide, osmium is rarely used in its pure state, and is instead often alloyed with other metals. Those alloys are utilized in high-wear applications. Osmium alloys such as osmiridium are very hard and, along with other platinum group metals, are used in the tips of fountain pens, instrument pivots, and electrical contacts, as they can resist wear from frequent operation. The stylus (needle) in early phonograph designs was also made of osmium, especially for 78-rpm records, until sapphire and synthetic diamond replaced the metal in later designs for 45-rpm and 33-rpm long-playing records.[42]
Electron micrograph of (organic) plant tissue without (top) and with (bottom) OsO4 staining
The Sharpless dihydroxylation: RL = Largest substituent; RM = Medium-sized substituent; RS = Smallest substituent
Osmium tetroxide has been used in fingerprint detection[43] and in staining fatty tissue for optical and electron microscopy. As a strong oxidant, it cross-links lipids mainly by reacting with unsaturated carbon-carbon bonds, and thereby both fixes biological membranes in
Osmium
8
place in tissue samples and simultaneously stains them. Because osmium atoms are extremely electron dense, osmium staining greatly enhances image contrast in transmission electron microscopy (TEM) studies of biological materials. Those carbon materials have otherwise very weak TEM contrast (see image).[14] Another osmium compound, osmium [44] ferricyanide (OsFeCN), exhibits similar fixing and staining action. An alloy of 90% platinum and 10% osmium is used in surgical implants such as pacemakers and replacement of pulmonary valves.[45] The tetroxide and a related compound, potassium osmate, are important oxidants for chemical synthesis, despite being very poisonous. For the Sharpless asymmetric dihydroxylation, which uses osmate for the conversion of a double bond into a vicinal diol, Karl Barry Sharpless won the Nobel Prize in Chemistry in 2001.[46] [47] In 1898 an Austrian chemist, Auer von Welsbach, developed the Oslamp with a filament made of osmium, which he introduced commercially in 1902. After only few years, osmium was replaced by the more stable metal tungsten (also known as wolfram). Tungsten has the highest melting point of any metal, and using it in light bulbs increases the luminous efficacy and life of incandescent lamps.[26] The light bulb manufacturer OSRAM (founded in 1906 when three German companies, Auer-Gesellschaft, AEG and Siemens & Halske, combined their lamp production facilities) derived its name from the elements of OSmium and wolfRAM.[50] Like
Post-flight appearance of Os, Ag, and Au mirrors from the front (left images) and rear panels of Space Shuttle. Blackening reveals oxidation due to irradiation by oxygen [48] [49] atoms.
palladium,
powdered
osmium
effectively
absorbs
hydrogen atoms. This could make osmium a potential candidate for a metal hydride battery electrode. However, osmium is expensive and would react with potassium hydroxide, the most common battery electrolyte.[51] Osmium has high reflectivity in the ultraviolet range of the
electromagnetic spectrum; for example, at 600 Å osmium has a reflectivity two times that of gold.[52] This high reflectivity is desirable in space-based UV spectrometers which have reduced mirror sizes due to space limitations. Osmium-coated mirrors were flown in several space missions aboard the Space Shuttle, but it soon became clear that the oxygen radicals in the low earth orbit are abundant enough to significantly deteriorate the osmium layer.[53]
Precautions Finely divided metallic osmium is pyrophoric.[41] Osmium reacts with oxygen at room temperature forming volatile osmium tetroxide. Some osmium compounds are also converted to the tetroxide if oxygen is present.[] This makes osmium tetroxide the main source for the contact to the environment. Osmium tetroxide is highly volatile and penetrates skin readily, and is very toxic by inhalation, ingestion, and skin contact.[54] Airborne low concentrations of osmium tetroxide vapor can cause lung congestion and skin or eye damage, and should therefore be used in a fume hood.[10] Osmium tetroxide is rapidly reduced to relatively inert compounds by polyunsaturated vegetable oils, such as corn oil.[55]
Osmium
9
External links • WebElements.com: Osmium
[56]
pnb:میمسوا
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Article Sources and Contributors
Article Sources and Contributors Osmium Source: http://en.wikipedia.org/w/index.php?oldid=309143033 Contributors: Abrech, Achim1999, Ahoerstemeier, AlimanRuna, AllStarZ, Allstarecho, Antandrus, Archimerged, Arkuat, AssegaiAli, Ataru, Atm1994, Auntof6, Baccyak4H, Batmanand, Bearcat, Bearings, Beetstra, Belovedfreak, Benjah-bmm27, Benjiboi, BlueEarth, Bobo192, Bomac, Borislav Dopudja, Brian Huffman, Bryan Derksen, Bucephalus, CYD, Can't sleep, clown will eat me, Canthusus, Capricorn42, Carnildo, Catbar, Chancellarius, CharonX, Cjtalamo, Conversion script, Cryptic C62, Cureden, Dadamonz, David Latapie, Dendodge, Derek Ross, DocWatson42, Donarreiskoffer, Dwmyers, Ed Nieuwenhuys, Edgar181, Edward321, Emperorbma, Eog1916, Eric-Wester, Erik Zachte, Erik9, Femto, Finlay McWalter, Firsfron, Foobar, FourteenDays, Fvw, Gaius Cornelius, Gbr3, George The Dragon, Giftlite, Glenn, GreenCrayon, GregorB, Grendelkhan, Herbee, Hu12, Icairns, Ideyal, ImperatorExercitus, JaGa, Jaan513, Jaraalbe, Jay Litman, Jdurg, Jeronimo, Jerzy, Jimp, Jklemens, Joanjoc, Joel Butts, Kastein, King of Hearts, Klparrot, KnowledgeOfSelf, Korath, Kwamikagami, LA2, Lavateraguy, Lehkost, Liftarn, Litherlandsand, Lockesdonkey, Marc Venot, Materialscientist, Mav, Mdf, Minesweeper, MoogleDan, Mr. Lefty, Mschel, Myasuda, Mygerardromance, Mysid, Nekura, Nergaal, Nihiltres, Odysseus1138, Oliver202, Omphacite, P3d0, PRiis, Pgk, Pip2andahalf, Plantsurfer, PlatinumX, Pollinator, Poolkris, Pras, Pwjb, RTC, Rebroad, Recognizance, RedWolf, Remember, Retiono Virginian, Reza kalani, Rhodiumking, Rifleman 82, Roberta F., RodC, Romanm, Rossnorman, Saperaud, Schneelocke, Sengkang, Sfuerst, Shaddack, Shanedidona, ShaunMacPherson, Shimmin, SirDanLucas, SkyLined, Sl, Sleigh, Snoyes, Squids and Chips, Stifynsemons, Stone, TUF-KAT, Tagishsimon, Tailpig, TantalumTelluride, Tavilis, Tetracube, Texture, The way, the truth, and the light, Thomas74, Thricecube, Thumperward, Tristanb, UltimateDestroyerOfWorlds, UtherSRG, V1adis1av, Van Slayer, Vsmith, Warut, Watch37264, WhatamIdoing, WuzUpYall, Yath, Yekrats, Yyy, Zahnrad, 272 anonymous edits
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12
Iridium
1
Iridium osmium ← iridium → platinum Rh ↑ Ir ↓ Mt Periodic Table Extended Periodic Table General Name, symbol, number
iridium, Ir, 77
Element category
transition metals
Group, Period, Block
9, 6, d
Appearance
silvery white
Standard atomic weight
−1
192.217(3) g·mol
Electron configuration
14
[Xe] 4f
Electrons per shell
7
2
5d 6s
2, 8, 18, 32, 15, 2 Physical properties
Phase
solid Density (near r.t.)
22.56 g·cm−3
Liquid density at m.p.
19 g·cm−3
Melting point
2739 K (2466 °C, 4471 °F) Boiling point
4701 K (4428 °C, 8002 °F)
Heat of fusion
41.12 kJ·mol−1
Heat of vaporization
563 kJ·mol−1
Specific heat capacity
(25 °C) 25.10 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
2713
2957
3252
3614
4069
4659
Iridium
2 Atomic properties Crystal structure
Oxidation states
face centered cubic −3,−1, 0, 1, 2, 3, 4, 5, 6
Electronegativity Ionization energies
2.20 (Pauling scale) 1st: 880 kJ/mol 2nd: 1600 kJ/mol
Atomic radius Covalent radius
136 pm 141±6 pm Miscellaneous [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(20 °C) 47.1 n Ω·m
Thermal conductivity
(300 K) 147 W·m
Thermal expansion
(25 °C) 6.4 µm·m
Speed of sound (thin rod)
(20 °C) 4825 m/s
Young's modulus
528 GPa
Shear modulus
210 GPa Bulk modulus
−1
−1
320 GPa
Poisson ratio
0.26
Mohs hardness
6.5
Vickers hardness Brinell hardness CAS registry number
−1
·K
1760 MPa 1670 MPa 7439-88-5 Most-stable isotopes
−1
·K
Iridium
3
Main article: Isotopes of iridium iso
NA
half-life
DM
DE (MeV)
DP
188
syn
1.73 d
ε
1.64
188
189
syn
13.2 d
ε
0.532
189
190
syn
11.8 d
ε
2.000
190
β-
1.460
192
ε
1.046
192
IT
0.161
192
IT
0.080
193
β
-
2.247
194
IT
?
194
Ir Ir Ir
37.3%
191
Ir
syn
192
Ir
192m2
Ir
syn 62.7%
193
Ir
193m
syn
194
syn
Ir
Ir
194m2
Ir
syn
Os Os Os
191
Ir is stable with 114 neutron
73.827 d
241 y
Pt Os Ir
193
Ir is stable with 116 neutron
10.5 d 19.3 h 171 d
Ir Pt Ir
References
Iridium (pronounced /ɨˈrɪdiəm/) is the chemical element with atomic number 77, and is represented by the symbol Ir. A very hard, brittle, silvery-white transition metal of the platinum family, iridium is the second densest element (after osmium) and is the most corrosion-resistant metal, even at temperatures as high as 2000 °C. Although only certain molten salts and halogens are corrosive to solid iridium, finely divided iridium dust is much more reactive and can even be flammable. The most important iridium compounds in terms of use are the salts and acids it forms with chlorine, though iridium also forms a number of 191 193 organometallic compounds used in catalysis and in research. Ir and Ir are the only two naturally occurring isotopes of iridium as well as the only stable isotopes; the latter is the more abundant of the two. Iridium was discovered in 1803 by Smithson Tennant in London, England, among insoluble impurities in natural platinum from South America. Although it is one of the rarest elements in the Earth's crust, with annual production and consumption of only three tonnes, it has a number of specialized industrial and scientific applications. Iridium is employed when high corrosion resistance and high temperatures are needed, as in spark plugs, crucibles for recrystallization of semiconductors at high temperatures, electrodes for the production of chlorine in the chloralkali process, and radioisotope thermoelectric generators used in unmanned spacecraft. Iridium compounds also find applications as catalysts for the production of acetic acid. An unusually high abundance of iridium in a clay layer of the K–T geologic boundary was a crucial clue that led to the theory that the extinction of dinosaurs and many other species 65 million years ago was caused by the impact of a massive extraterrestrial object—the so-called Alvarez hypothesis. Iridium is found in meteorites with an abundance much higher than its average abundance in the Earth's crust. It is thought that the amount of iridium in
Iridium
4
the planet Earth is much higher than what is observed in crustal rocks, but because of the high density and tendency of iridium to bond with iron, most iridium descended below the crust and into the Earth's core when the planet was young and still molten.
Characteristics Physical
1 troy ounce (31 g) of arc-melted iridium
A member of the platinum group metals, iridium is white, resembling platinum, but with a slight yellowish cast. Due to its hardness, brittleness, and very high melting point (the tenth highest of all elements), solid iridium is difficult to machine, form, or work, and thus powder metallurgy is commonly employed instead.[2] It is the only metal to maintain good mechanical properties in air at temperatures above 1600 °C.[3] Iridium has a very high boiling point (11th among all elements) and becomes a superconductor under 0.14 K.[4]
Iridium's modulus of elasticity is the second highest among the metals, only being surpassed by osmium.[3] This, together with a high modulus of rigidity and a very low figure for Poisson's ratio (the relationship of longitudinal to lateral strain), indicate the high degree of stiffness and resistance to deformation that have rendered its fabrication into useful components a matter of great difficulty. Despite these limitations and iridium's high cost, a number of applications have developed where mechanical strength is an essential factor in some of the extremely severe conditions encountered in modern technology.[3] The measured density of iridium is only slightly lower (by about 0.1%) than that of osmium, the densest element known.[5] [6] There had been some ambiguity regarding which of the two element was denser, due to the small size of the difference in density and difficulties in measuring it accurately,[7] but, with increased accuracy in factors used for calculating density X-ray crystallographic data yielded densities of 22.56 g/cm3 for iridium and 22.59 g/cm3 for osmium.[8]
Chemical Iridium is the most corrosion-resistant metal known:[9] it is not attacked by any acid, by aqua regia, by any molten metals, or by silicates at high temperatures. It can, however, be attacked by some molten salts, such as sodium cyanide and potassium cyanide,[9] as well as oxygen and the halogens (particularly fluorine)[10] at higher temperatures.[11]
Compounds Oxidation states [12] of iridium −3
[Ir(CO)3]3−
−1
[Ir(CO)3(PPh3)]−
Iridium
5
0
Ir4(CO)12
+1
[Ir(CO)Cl(PPh3)2]
+2
IrCl2
+3
IrCl3
+4
IrO2
+5
Ir4F20
+6
IrF6
Iridium forms compounds in oxidation states between −3 to +6; the most common oxidation states are +3 and +4.[2] Well-characterized examples of the highest oxidation state are rare, but include IrF6 and two mixed oxides Sr2MgIrO6 and Sr2CaIrO6.[2] [13]
Iridium dioxide, IrO2, a brown powder, is the only well-characterized oxide of iridium.[2] A sesquioxide, Ir2O3, has been described as a blue-black powder which is oxidized to IrO2 by HNO3.[10] The corresponding disulfides, diselenides, sesquisulfides and sesquiselenides are known and IrS3 has also been reported.[2] Iridium also forms iridates with oxidation states +4 and +5, such as K2IrO3 and KIrO3, which can be prepared from the reaction of potassium oxide or potassium superoxide with iridium at high temperatures.[14] While no binary hydrides of iridium, IrxHy are known, complexes are known that contain IrH4−5 and IrH3−6, where iridium has the +1 and +3 oxidation states, respectively.[15] The ternary hydride Mg6Ir2H11 is believed to contain both the IrH4−5 and the 18-electron IrH5−4 anion.[16] No monohalides or dihalides are known, whereas trihalides, IrX3, are known for all of the halogens.[2] For oxidation states +4 and above, only the tetrafluoride, pentafluoride and hexafluoride are known.[2] Iridium hexafluoride, IrF6, is a volatile and highly reactive yellow solid, composed of octahedral molecules. It decomposes in water and is reduced to IrF4, a crystalline solid, by iridium black.[2] Iridium pentafluoride has similar properties but it is actually a tetramer, Ir4F20, formed by four corner-sharing octahedra.[2] Hexachloroiridic(IV) acid, H IrCl , and its ammonium 2
6
salt are the most important iridium compounds from an industrial perspective.[17] They are involved in the purification of iridium and used as precursors for most other iridium compounds, as well as in the preparation of anode coatings. The IrCl2−6 ion has an intense dark brown color, and can be readily reduced to the Vaska's complex lighter-colored IrCl3−6 and vice versa.[17] Iridium trichloride, IrCl3, which can be obtained in anhydrous form from direct oxidation of iridium powder by chlorine at 650 °C,[17] or in hydrated form by dissolving Ir2O3 in hydrochloric acid, is often used as a starting material for the synthesis of other Ir(III) compounds.[2] Another compound used as a starting material is ammonium hexachloroiridate(III), (NH4)3IrCl6. Iridium(III) complexes are diamagnetic (low-spin) and generally have an octahedral molecular geometry.[2] Organoiridium compounds contain iridium–carbon bonds where the metal is usually in lower oxidation states. For example, oxidation state zero is found in tetrairidium dodecacarbonyl, Ir4(CO)12, which is the most common and stable binary carbonyl of iridium.[2] In this compound, each of the iridium atoms is bonded to the other three,
Iridium forming a tetrahedral cluster. Some organometallic Ir(I) compounds are notable enough to be named after their discoverers. One is Vaska's complex, IrCl(CO)[P(C6H5)3]2, which has the unusual property of binding to the dioxygen molecule, O2.[18] Another one is Crabtree's catalyst, a homogeneous catalyst for hydrogenation reactions.[19] These compounds are 8 both square planar, d complexes, with a total of 16 valence electrons, which accounts for their reactivity.[20]
Isotopes Iridium has two naturally occurring, stable isotopes, 191Ir and 193Ir, with natural abundances of 37.3% and 62.7%, respectively.[21] At least 34 radioisotopes have also been synthesized, ranging in mass number from 164 to 199. Twenty-seven of these are lighter than the stable isotopes, while six are heavier. 192Ir, which falls between the two stable isotopes, is the most stable radioisotope, with a half-life of 73.827 days, and finds application in brachytherapy.[22] Three other isotopes have half-lives of at least a day—188Ir, 189Ir, 190Ir.[21] One of the least stable isotopes is 165Ir with a half-life of 1 µs. Isotopes with masses below 191 decay by some combination of β+ decay, α decay, and proton emission, with the exceptions of 189Ir, which decays by electron capture, and 190Ir, which decays by positron emission. Synthetic isotopes heavier than 191 decay by β− decay, although 192Ir also has a minor electron capture decay path.[21] All known isotopes of iridium were discovered between 1934 and 2001; the most recent is 171Ir.[23] At least 32 metastable isomers have been characterized, ranging in mass number from 164 to 197. The most stable of these is 192m2Ir, which decays by isomeric transition with a half-life of 241 years,[21] making it more stable than any of iridium's synthetic isotopes in their ground states. The least stable isomer is 190m3Ir with a half-life of only 2 µs.[21] The isotope 191Ir was the first one of any element to be shown to present a Mössbauer effect. This renders it useful for Mössbauer spectroscopy for research in physics, chemistry, biochemistry, metallurgy, and mineralogy.[24]
History The discovery of iridium is intertwined with that of platinum and the other metals of the platinum group. Native platinum used by ancient Ethiopians[25] and by South American cultures[26] always contained a small amount of the other platinum group metals, including iridium. Platinum reached Europe as platina ("small silver"), found in the 17th century by the Spanish conquerors in a region today known as the department of Chocó in Colombia.[27] The discovery that this metal was not an alloy of known elements, but instead a distinct new element, did not occur until 1748.[28] Chemists who studied platinum dissolved it in aqua regia (a mixture of hydrochloric and nitric acids) to create soluble salts. They always observed a small amount of a dark, insoluble residue.[3] Joseph Louis Proust thought that the residue was graphite.[3] The French chemists Victor Collet-Descotils, Antoine François, comte de Fourcroy, and Louis Nicolas Vauquelin also observed the black residue in 1803, but did not obtain enough for further experiments.[3] In 1803, British scientist Smithson Tennant (1761–1815) analyzed the insoluble residue and concluded that it must contain a new metal. Vauquelin treated the powder alternatively with alkali and acids[9] and obtained a volatile new oxide, which he believed to be of this new metal—which he named ptene, from the Greek word πτηνος (ptènos) for winged.[29]
6
Iridium [30]
Tennant, who had the advantage of a much greater amount of residue, continued his research and identified the two previously undiscovered elements in the black residue, iridium and osmium.[3] [9] He obtained dark red crystals (probably of Na2[IrCl6]·nH2O) by a sequence of reactions with sodium hydroxide and hydrochloric acid.[30] He named iridium after Iris (Ιρις), the Greek winged goddess of the rainbow and the messenger of the Olympian gods, because many of the salts he obtained were strongly colored.[31] [32] Discovery of the new elements was documented in a letter to the Royal Society on June 21, 1804.[3] [33] British scientist John George Children was the first to melt a sample of iridium in 1813 with the aid of "the greatest galvanic battery that has ever been constructed" (at that time).[3] The first to obtain high purity iridium was Robert Hare in 1842. He found that it had a density of around 21.8 g/cm3 and noted that the metal is nearly unmalleable and very hard. The first melting in appreciable quantity was done by Henri Sainte-Claire Deville and Jules Henri Debray in 1860. They required burning more than 300 L of pure O2 and H2 for each kilogram of iridium.[3] These extreme difficulties in melting the metal limited the possibilities for handling iridium. John Isaac Hawkins was looking to obtain a fine and hard point for fountain pen nibs The Greek goddess Iris, after and in 1834 managed to create an iridium-pointed gold pen. whom Iridium was named. In 1880 John Holland and William Lofland Dudley were able to melt iridium by adding phosphorus and patented the process in the United States; British company Johnson Matthey later stated that they had been using a similar process since 1837 and had already presented fused iridium at a number of World Fairs.[3] The first use of an alloy of iridium with ruthenium in thermocouples was made by Otto Feussner in 1933. These allowed for the measurement of high temperatures in air up to 2000 °C.[3] In 1957 Rudolf Mössbauer, in what has been called one of the "landmark experiments in twentieth century physics",[34] discovered the resonant and recoil-free emission and absorption of gamma rays by atoms in a solid metal sample containing only 191Ir.[35] This phenomenon, known as the Mössbauer effect (which has since been observed for other nuclei, such as 57Fe), and developed as Mössbauer spectroscopy, has made important contributions to research in physics, chemistry, biochemistry, metallurgy, and mineralogy.[24] Mössbauer received the Nobel Prize in Physics in 1961, just three years after he published his discovery.[36]
7
Iridium
8
Occurrence Iridium is one of the least abundant elements in the Earth's crust, having an average mass fraction of 0.001 ppm in crustal rock; gold is 4 times more abundant, platinum is 10 times more abundant, and silver and mercury are 80 times more abundant.[2] Tellurium is about as abundant as iridium, and only three naturally occurring elements are less abundant: rhenium, ruthenium, and rhodium, iridium being 10 times more abundant than the last two.[2] In contrast to its low abundance in crustal rock, iridium is relatively common in meteorites, with concentrations of 0.5 ppm or more.[38] It is thought that the overall concentration of iridium on Earth is much higher than what is observed in crustal rocks, but because of the density and siderophilic ("iron-loving") character of iridium, it descended below the crust and into the Earth's core when the planet was still molten.[17]
The Willamette Meteorite, the sixth largest meteorite found in the world, [37] has 4.7 ppm iridium.
Iridium is found in nature as an uncombined element or in natural alloys; especially the iridium–osmium alloys, osmiridium (osmium rich), and iridiosmium (iridium rich).[9] In the nickel and copper deposits the platinum group metals occur as sulfides (i.e. (Pt,Pd)S)), tellurides (i.e. PtBiTe), antimonides (PdSb), and arsenides (i.e. PtAs2). In all of these compounds platinum is exchanged by a small amount of iridium and osmium. As with all of the platinum group metals, iridium can be found naturally in alloys with raw nickel or raw copper.[39] Within the Earth's crust, iridium is found at highest concentrations in three types of geologic structure: igneous deposits (crustal intrusions from below), impact craters, and deposits reworked from one of the former structures. The largest known primary reserves [40] are in the Bushveld igneous complex in South Africa, though the large copper–nickel deposits near Norilsk in Russia, and the Sudbury Basin in Canada are also significant sources of iridium. Smaller reserves are found in the United States.[40] Iridium is also found in secondary deposits, combined with platinum and other platinum group metals in alluvial deposits. The alluvial deposits used by pre-Columbian people in the Chocó Department of Colombia are still a source for platinum-group metals. As of 2003 the world reserves had not been estimated.[9]
Iridium
9
K–T boundary presence The K–T boundary of 65 million years ago, marking the temporal border between the Cretaceous and Tertiary periods of geological time, was identified by a thin stratum of iridium-rich clay.[41] A team led by Luis Alvarez proposed in 1980 an extraterrestrial origin for this iridium, attributing it to an asteroid or comet impact.[41] Their theory, known as the Alvarez hypothesis, is now widely accepted to explain the The red arrow points to the K–T demise of the dinosaurs. A large buried impact crater boundary. structure with an estimated age of about 65 million years was later identified under what is now the Yucatán Peninsula (the Chicxulub [42] [43] crater). Dewey M. McLean and others argue that the iridium may have been of volcanic origin instead, as the Earth's core is rich in iridium, and active volcanoes such as Piton de la Fournaise, in the island of Réunion, are still releasing iridium.[44] [45]
Production Year
Price [46] [47] ($/ozt)
2001
415.25
2002
294.62
2003
93.02
2004
185.33
2005
169.51
2006
349.45
2007
440.00
Iridium is obtained commercially as a by-product from nickel and copper mining and processing. During electrorefining of copper and nickel, noble metals such as silver, gold and the platinum group metals as well as selenium and tellurium settle to the bottom of the cell as anode mud, which forms the starting point for their extraction.[46] [48] In order to separate the metals, they must first be brought into solution. Several methods are available depending on the separation process and the composition of the mixture; two representative methods are fusion with sodium peroxide followed by dissolution in aqua regia, and dissolution in a mixture of chlorine with hydrochloric acid.[17] [40] After it is dissolved, iridium is separated from the other platinum group metals by precipitating (NH4)2IrCl6 or by extracting IrCl2−6 with organic amines.[49] The first method is similar to the procedure Tennant and Wollastone used for their separation. The second method can be planned as continuous liquid–liquid extraction and is therefore more suitable for industrial scale production. In either case, the product is reduced using hydrogen, yielding the metal as a powder or sponge that can be treated using powder metallurgy techniques.[50] [51] Annual production of iridium circa 2000 was around 3 tonnes or about 100,000 troy ounces (ozt).[52] [9] The price of iridium as of 2007 was $440 USD/ozt,[46] but the price fluctuates
Iridium considerably, as shown in the table. The high volatility of the prices of the platinum group metals has been attributed to supply, demand, speculation, and hoarding, amplified by the small size of the market and instability in the producing countries.[53]
Applications The global demand for iridium in 2007 was 119,000 troy ounces (3,700 kg), out of which 25,000 ozt (780 kg) were used for electrical applications such as spark plugs; 34,000 ozt (1,100 kg) for electrochemical applications such as electrodes for the chloralkali process; 24,000 ozt (750 kg) for catalysis; and 36,000 ozt (1,100 kg) for other uses.[54]
Industrial and medical The high melting point, hardness and corrosion resistance of iridium and its alloys determine most of its applications. Iridium and especially iridium–platinum alloys or osmium–iridium alloys have a low wear and are used, for example, for multi-pored spinnerets, through which a plastic polymer melt is extruded to form fibers, such as rayon.[55] Osmium–iridium is used for compass bearings and for balances.[9] Corrosion and heat resistance makes iridium an important alloying agent. Certain long-life aircraft engine parts are made of an iridium alloy and an iridium–titanium alloy is Molecular structure of Ir(mppy)3 used for deep-water pipes because of its corrosion resistance.[9] Iridium is also used as a hardening agent in platinum alloys. The Vickers hardness of pure platinum is 56 HV while platinum with 50% of iridium can reach over 500 HV.[56] [57] Devices that must withstand extremely high temperatures are often made from iridium. For example, high-temperature crucibles made of iridium are used in the Czochralski process to produce oxide single-crystals (such as sapphires) for use in computer memory devices and in solid state lasers.[58] [59] The crystals, such as gadolinium gallium garnet and yttrium gallium garnet, are grown by melting pre-sintered charges of mixed oxides under oxidizing conditions at temperatures up to 2100 °C.[3] Its resistance to arc erosion makes iridium alloys ideal for electrical contacts for spark plugs.[59] [60] Iridium compounds are used as catalysts in the Cativa process for carbonylation of methanol to produce acetic acid.[61] Iridium itself is used as a catalyst in a type of automobile engine introduced in 1996 called the direct-ignition engine.[9] The radioisotope iridium-192 is one of the two most important sources of energy for use in industrial γ-radiography for non-destructive testing of metals.[62] [63] Additionally, 192Ir is used as a source of gamma radiation for the treatment of cancer using brachytherapy, a form of radiotherapy where a sealed radioactive source is placed inside or next to the area requiring treatment. Specific treatments include high dose rate prostate brachytherapy, bilary duct brachytherapy, and intracavitary cervix brachytherapy.[9]
10
Iridium
11
Scientific An alloy of 90% platinum and 10% iridium was used in 1889 to construct the International Prototype Meter and kilogram mass, kept by the International Bureau of Weights and Measures near Paris.[9] The meter bar was replaced as the definition of the fundamental unit of length in 1960 by a line in the atomic spectrum of krypton,[64] [65] but the kilogram prototype is still the international standard of mass.[66]
International Prototype Meter bar
Iridium has been used in the radioisotope thermoelectric generators of unmanned spacecraft such as the Voyager, Viking, Pioneer, Cassini, Galileo, and New Horizons. Iridium was chosen to encapsulate the plutonium-238 fuel in the generator because it can withstand the operating temperatures of up to 2000 °C and for its great strength.[3] Another use concerns X-ray optics, especially X-ray telescopes.[67] The mirrors of the Chandra X-ray Observatory are coated with a layer of iridium 60 nm thick. Iridium proved to be the best choice for reflecting X-rays after nickel, gold, and platinum were tested. The iridium layer, which had to be smooth to within a few atoms, was applied by depositing iridium vapor under high vacuum on a base layer of chromium.[68] Iridium is used in particle physics for the production of antiprotons, a form of antimatter. Antiprotons are made by shooting a high-intensity proton beam at a conversion target, which needs to be made from a very high density material. Although tungsten may be used instead, iridium has the advantage of better stability under the shock waves induced by the temperature rise due to the incident beam.[69] Carbon–hydrogen
bond
activation (C–H activation) is an area of research on reactions that cleave carbon–hydrogen bonds, [70] which were traditionally Oxidative addition to hydrocarbons in organoiridium chemistry. [71] regarded as unreactive. The first reported successes at activating C–H bonds in saturated hydrocarbons, published in 1982, used organometallic iridium complexes that undergo an oxidative addition with the hydrocarbon.[70] [71] Iridium complexes are being investigated as catalysts for asymmetric hydrogenation. These catalysts have been used in the synthesis of natural products and able to hydrogenate certain difficult substrates, such as unfunctionalized alkenes, enantioselectively (generating only one of the two possible enantiomers).[72] [73] Iridium forms a variety of complexes of fundamental interest in triplet harvesting.[74] [76]
[75]
Iridium
12
Historical Iridium–osmium alloys were used to tip fountain pen nibs. The first major use of iridium was in 1834 in nibs mounted on gold.[3] Since 1944, the famous Parker 51 fountain pen was fitted with a nib tipped by a ruthenium and iridium alloy (with 3.8% iridium). The tip material in modern fountain pens is still conventionally called "iridium," although there is seldom any iridium in it; other metals such as tungsten have taken its place.[77]
Fountain pen nib labeled Iridium Point
An iridium–platinum alloy was used for the touch holes or vent pieces of cannons. According to a report of the Paris Exhibition of 1867, one of the pieces being exhibited by Johnson and Matthey "has been used in a Withworth gun for more than 3000 rounds, and scarcely shows signs of wear yet. Those who know the constant trouble and expense which are occasioned by the wearing of the vent-pieces of cannon when in active service, will appreciate this important adaptation".[78] The pigment iridium black, which consists of very finely divided iridium, is used for painting porcelain an intense black; it was said that "all other porcelain black colors appear grey by the side of it".[79]
Precautions Iridium in bulk metallic form is not biologically important or hazardous to health due to its lack of reactivity with tissues; there are only about 20 parts per trillion of iridium in human tissue.[9] However, finely divided iridium powder can be hazardous to handle, as it is an irritant and may ignite in air.[40] Very little is known about the toxicity of iridium compounds because they are used in very small amounts, but soluble salts, such as the iridium halides, could be hazardous due to elements other than iridium or due to iridium itself.[22] However, most iridium compounds are insoluble, which makes absorption into the body difficult.[9] A radioisotope of iridium, 192Ir, is dangerous like other radioactive isotopes. The only reported injuries related to iridium concern accidental exposure to radiation from 192Ir used in brachytherapy.[22] High-energy gamma radiation from 192Ir can increase the risk of cancer. External exposure can cause burns, radiation poisoning, and death. Ingestion of 192 Ir can burn the linings of the stomach and the intestines.[80] 192Ir, 192mIr, and 194mIr tend to deposit in the liver, and can pose health hazards from both gamma and beta radiation.[38]
Iridium
See also • Centers for Disease Control and Prevention
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. pp. 1113–1143,1294. ISBN 0-7506-3365-4. OCLC 213025882 37499934 41901113 (http:/ / worldcat. org/ oclc/ 213025882+ 37499934+ 41901113). [3] Hunt, L. B. (1987). " A History of Iridium (http:/ / www. platinummetalsreview. com/ dynamic/ article/ view/ pmr-v31-i1-032-041)". Platinum Metals Review 31 (1): 32–41. . [4] Kittel, C. (2004). Introduction to Solid state Physics, 7th Edition. Wiley-India. ISBN 8126510455. [5] Arblaster, J. W. (1995). " Osmium, the Densest Metal Known (http:/ / www. platinummetalsreview. com/ dynamic/ article/ view/ pmr-v39-i4-164-164)". Platinum Metals Review 39 (4): 164. . [6] Cotton, Simon (1997). Chemistry of Precious Metals. Springer-Verlag New York, LLC. p. 78. [7] Lide, D. R. (1990). CRC Handbook of Chemistry and Physics (70th Edn.). Boca Raton (FL):CRC Press. [8] Arblaster, J. W. (1989). " Densities of osmium and iridium: recalculations based upon a review of the latest crystallographic data (http:/ / www. platinummetalsreview. com/ pdf/ pmr-v33-i1-014-016. pdf)" (PDF). Platinum Metals Review 33 (1): 14–16. . [9] Emsley, J. (2003). "Iridium". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. pp. 201–204. ISBN 0198503407. [10] Perry, D. L. (1995). Handbook of Inorganic Compounds. CRC Press. pp. 203–204. ISBN 0-8492-8671-3. [11] Lagowski, J. J., ed (2004). Chemistry Foundations and Applications. 2. Thomson Gale. pp. 250–251. ISBN 0-02-865732-3. [12] Common oxidation states are in bold. [13] Jung, D. (1995). "High Oxygen Pressure and the Preparation of New Iridium (VI) Oxides with Perovskite Structure: Sr2MIrO6 (M = Ca, Mg)". Journal of Solid State Chemistry 115 (2): 447–455. doi: 10.1006/jssc.1995.1158 (http:/ / dx. doi. org/ 10. 1006/ jssc. 1995. 1158). [14] Gulliver, D. J; Levason, W. (1982). "The chemistry of ruthenium, osmium, rhodium, iridium, palladium and platinum in the higher oxidation states". Coordination Chemistry Reviews 46: 1–127. doi: 10.1016/0010-8545(82)85001-7 (http:/ / dx. doi. org/ 10. 1016/ 0010-8545(82)85001-7). [15] Holleman, A. F.; Wiberg, E.; Wiberg, N. (2001). Inorganic Chemistry, 1st Edition. Academic Press. ISBN 0123526515. OCLC 47901436 (http:/ / worldcat. org/ oclc/ 47901436). [16] Černý, R.; Joubert, J.-M.; Kohlmann, H.; Yvon, K. (2002). "Mg6Ir2H11, a new metal hydride containing saddle-like IrH5−4 and square-pyramidal IrH4−5 hydrido complexes". Journal of Alloys and Compounds 340 (1-2): 180–188. doi: 10.1016/S0925-8388(02)00050-6 (http:/ / dx. doi. org/ 10. 1016/ S0925-8388(02)00050-6). [17] Renner, H.; Schlamp, G.; Kleinwächter, I.; Drost, E.; Lüschow, H. M.; Tews, P.; Panster, P.; Diehl, M.; Lang, J.; Kreuzer, T.; Knödler, A.; Starz, K. A.; Dermann, K.; Rothaut, J.; Drieselman, R. (2002). "Platinum group metals and compounds". Ullmann's Encyclopedia of Industrial Chemistry. Wiley. doi: 10.1002/14356007.a21_075 (http:/ / dx. doi. org/ 10. 1002/ 14356007. a21_075). [18] Vaska, L.; DiLuzio, J.W. (1961). "Carbonyl and Hydrido-Carbonyl Complexes of Iridium by Reaction with Alcohols. Hydrido Complexes by Reaction with Acid". Journal of the American Chemical Society 83: 2784–2785. doi: 10.1021/ja01473a054 (http:/ / dx. doi. org/ 10. 1021/ ja01473a054). [19] Crabtree, R. H. (1979). "Iridium compounds in catalysis". Accounts of Chemical Research 12: 331–337. doi: 10.1021/ar50141a005 (http:/ / dx. doi. org/ 10. 1021/ ar50141a005). [20] Crabtree, R. H. (2005). The Organometallic Chemistry of the Transition Metals (http:/ / www. wiley. com/ WileyCDA/ WileyTitle/ productCd-0471662569. html). Wiley. ISBN 978-0-471-66256-3. OCLC 224478241 (http:/ / worldcat. org/ oclc/ 224478241). . [21] Audi, G. (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3–128. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [22] Mager Stellman, J. (1998). " Iridium (http:/ / books. google. com/ books?id=nDhpLa1rl44C& pg=PT125)". Encyclopaedia of Occupational Health and Safety. International Labour Organization. pp. 63.19. ISBN 9789221098164. OCLC 35279504 45066560 (http:/ / worldcat. org/ oclc/ 35279504+ 45066560). . [23] Arblaster, J. W. (2003). " The discoverers of the iridium isotopes: the thirty-six known iridium isotopes found between 1934 and 2001 (http:/ / www. platinummetalsreview. com/ dynamic/ article/ view/ 47-4-167-174)". Platinum Metals Review 47 (4): 167–174. .
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Iridium [24] Chereminisoff, N. P. (1990). Handbook of Ceramics and Composites. CRC Press. p. 424. ISBN 082478006X. [25] Ogden, J. M. (1976). " The So-Called 'Platinum' Inclusions in Egyptian Goldwork (http:/ / www. jstor. org/ stable/ 3856354)". The Journal of Egyptian Archaeology 62: 138–144. doi: 10.2307/3856354 (http:/ / dx. doi. org/ 10. 2307/ 3856354). . [26] Chaston, J. C. (1980). "The Powder Metallurgy of Platinum". Platinum Metals Rev. 24 (21): 70–79. [27] McDonald, M. (959). " The Platinum of New Granada: Mining and Metallurgy in the Spanish Colonial Empire (http:/ / www. platinummetalsreview. com/ dynamic/ article/ view/ pmr-v3-i4-140-145)". Platinum Metals Review 3 (4): 140–145. . [28] Juan, J.; de Ulloa, A. (1748) (in Spanish). Relación histórica del viage a la América Meridional. 1. p. 606. [29] Thomson, T. (1831). A System of Chemistry of Inorganic Bodies. Baldwin & Cradock, London; and William Blackwood, Edinburgh. p. 693. [30] Griffith, W. P. (2004). "Bicentenary of Four Platinum Group Metals. Part II: Osmium and iridium – events surrounding their discoveries". Platinum Metals Review 48 (4): 182–189. doi: 10.1595/147106704X4844 (http:/ / dx. doi. org/ 10. 1595/ 147106704X4844). [31] Iridium literally means "of rainbows". [32] Weeks, M. E. (1968). Discovery of the Elements (7 ed.). Journal of Chemical Education. pp. 414–418. ISBN 0848685792. OCLC 23991202 (http:/ / worldcat. org/ oclc/ 23991202). [33] Tennant, S. (1804). " On Two Metals, Found in the Black Powder Remaining after the Solution of Platina (http:/ / www. jstor. org/ pss/ 107152)". Philosophical Transactions of the Royal Society of London 94: 411–418. doi: 10.1098/rstl.1804.0018 (http:/ / dx. doi. org/ 10. 1098/ rstl. 1804. 0018). . [34] Trigg, G. L. (1995). Landmark Experiments in Twentieth Century Physics. Courier Dover Publications. pp. 179–190. ISBN 048628526X. OCLC 31409781 (http:/ / worldcat. org/ oclc/ 31409781). [35] Mössbauer, R. L. (1958). "Gammastrahlung in Ir191" (in German). Zeitschrift für Physik a Hadrons and Nuclei 151: 124–143. doi: 10.1007/BF01344210 (http:/ / dx. doi. org/ 10. 1007/ BF01344210). [36] Waller, I. (1964). " The Nobel Prize in Physics 1961: presentation speech (http:/ / nobelprize. org/ nobel_prizes/ physics/ laureates/ 1961/ press. html)". Nobel Lectures, Physics 1942-1962. Elsevier. . [37] Scott, E. R. D.; Wasson, J. T.; Buchwald, V. F. (1973). "The chemical classification of iron meteorites—VII. A reinvestigation of irons with Ge concentrations between 25 and 80 ppm". Geochimica et Cosmochimica Acta 37: 1957–1983. doi: 10.1016/0016-7037(73)90151-8 (http:/ / dx. doi. org/ 10. 1016/ 0016-7037(73)90151-8). [38] " Iridium (http:/ / www. ead. anl. gov/ pub/ doc/ Iridium. pdf)" (PDF). Human Health Fact Sheet. Argonne National Laboratory. March 2005. . Retrieved 2008-09-20. [39] Xiao, Z.; Laplante, A. R. (2004). "Characterizing and recovering the platinum group minerals—a review". Minerals Engineering 17: 961–979. doi: 10.1016/j.mineng.2004.04.001 (http:/ / dx. doi. org/ 10. 1016/ j. mineng. 2004. 04. 001). [40] Seymour, R. J.; O'Farrelly, J. I. (2001). "Platinum-group metals". Kirk Othmer Encyclopedia of Chemical Technology. Wiley. doi: 10.1002/0471238961.1612012019052513.a01.pub2 (http:/ / dx. doi. org/ 10. 1002/ 0471238961. 1612012019052513. a01. pub2). [41] Alvarez, L. W.; Alvarez, W.; Asaro, F.; Michel, H. V. (1980). "Extraterrestrial cause for the Cretaceous–Tertiary extinction". Science 208 (4448): 1095–1108. doi: 10.1126/science.208.4448.1095 (http:/ / dx. doi. org/ 10. 1126/ science. 208. 4448. 1095). PMID 17783054. [42] Hildebrand, A. R.; Penfield, Glen T.; Kring, David A.; Pilkington, Mark; Zanoguera, Antonio Camargo; Jacobsen, Stein B.; Boynton, William V. (September 1991). " Chicxulub Crater; a possible Cretaceous/Tertiary boundary impact crater on the Yucatan Peninsula, Mexico (http:/ / geology. geoscienceworld. org/ cgi/ content/ abstract/ 19/ 9/ 867)". Geology 19 (9): 867–871. doi: 10.1130/0091-7613(1991)019<0867:CCAPCT>2.3.CO;2 (http:/ / dx. doi. org/ 10. 1130/ 0091-7613(1991)019<0867:CCAPCT>2. 3. CO;2). . [43] Frankel, C. (1999) (in English). The End of the Dinosaurs: Chicxulub Crater and Mass Extinctions. Cambridge University Press. ISBN 0521474477. OCLC 40298401 (http:/ / worldcat. org/ oclc/ 40298401). [44] Ryder, G.; Fastovsky, D. E.; Gartner, S. (1996). The Cretaceous-Tertiary Event and Other Catastrophes in Earth History. Geological Society of America. p. 47. ISBN 0813723078. [45] Toutain, J.-P.; Meyer, G. (1989). "Iridium-Bearing Sublimates at a Hot-Spot Volcano (Piton De La Fournaise, Indian Ocean)". Geophysical Research Letters 16 (12): 1391–1394. doi: 10.1029/GL016i012p01391 (http:/ / dx. doi. org/ 10. 1029/ GL016i012p01391). [46] George, M. W. (2008). " Platinum-group metals (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ platinum/ mcs-2008-plati. pdf)" (PDF). U.S. Geological Survey Mineral Commodity Summaries (USGS Mineral Resources Program). . [47] George, M. W. (2006). " Platinum-group metals (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ platinum/ platimcs06. pdf)" (PDF). U.S. Geological Survey Mineral Commodity Summaries (USGS Mineral Resources Program). .
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Iridium [48] George, M. W. (PDF). 2006 Minerals Yearbook: Platinum-Group Metals (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ platinum/ myb1-2006-plati. pdf). United States Geological Survey USGS. . [49] Gilchrist, Raleigh (1943). "The Platinum Metals.". Chemical Reviews 32 (3): 277–372. doi: 10.1021/cr60103a002 (http:/ / dx. doi. org/ 10. 1021/ cr60103a002). [50] Ohriner, E. K. (2008). "Processing of Iridium and Iridium Alloys". Platinum Metals Review 52 (3): 186–197. doi: 10.1595/147106708X333827 (http:/ / dx. doi. org/ 10. 1595/ 147106708X333827). [51] Hunt, L. B.; Lever, F. M. (1969). " Platinum Metals: A Survey of Productive Resources to industrial Uses (http:/ / www. platinummetalsreview. com/ pdf/ pmr-v13-i4-126-138. pdf)" (PDF). Platinum Metals Review 13 (4): 126–138. . [52] Like other precious metals, iridium is customarily traded in troy ounces, which are equivalent to about 31.1 grams. [53] Hagelüken, C. (2006). " Markets for the catalysts metals platinum, palladium, and rhodium (http:/ / www. preciousmetals. umicore. com/ publications/ articles_by_umicore/ general/ show_Metal_PGMmarkets_200602. pdf)" (PDF). Metall 60 (1–2): 31–42. . [54] Jollie, D. (May 2008) (PDF). Platinum 2008 (http:/ / www. platinum. matthey. com/ uploaded_files/ Pt2008/ 08_complete_publication. pdf). Johnson Matthey. . Retrieved 2008-10-13. [55] Egorova, R. V.; Korotkov, B. V.; Yaroshchuk, E. G.; Mirkus, K. A.; Dorofeev N. A.; Serkov, A. T. (1979). "Spinnerets for viscose rayon cord yarn". Fibre Chemistry 10 (4): 377–378. doi: 10.1007/BF00543390 (http:/ / dx. doi. org/ 10. 1007/ BF00543390). [56] Darling, A. S. (1960). " Iridium Platinum Alloys (http:/ / www. platinummetalsreview. com/ pdf/ pmr-v4-i1-018-026. pdf)" (PDF). Platinum Metals Review 4 (l): 18–26. . Retrieved 2008-10-13. [57] Biggs, T.; Taylor, S. S.; van der Lingen, E. (2005). "The Hardening of Platinum Alloys for Potential Jewellery Application". Platinum Metals Review 49 (1): 2–15. doi: 10.1595/147106705X24409 (http:/ / dx. doi. org/ 10. 1595/ 147106705X24409). [58] Crookes, W. (1908). " On the Use of Iridium Crucibles in Chemical Operations (http:/ / www. jstor. org/ pss/ 93031)". Proceedings of the Royal Society of London. Series A, Containing Papers of a Mathematical and Physical Character 80 (541): 535–536. doi: 10.1098/rspa.1908.0046 (http:/ / dx. doi. org/ 10. 1098/ rspa. 1908. 0046). . [59] Handley, J. R. (1986). " Increasing Applications for Iridium (http:/ / www. platinummetalsreview. com/ dynamic/ article/ view/ pmr-v30-i1-012-013)". Platinum Metals Review 30 (1): 12–13. . [60] Stallforth, H.; Revell, P. A. (2000). Euromat 99 (http:/ / books. google. de/ books?hl=de& lr=& id=I03qepnj2IwC). Wiley-VCH. . [61] Cheung, H.; Tanke, R. S.; Torrence, G. P. (2000). "Acetic acid". Ullmann's Encyclopedia of Industrial Chemistry. Wiley. doi: 10.1002/14356007.a01_045 (http:/ / dx. doi. org/ 10. 1002/ 14356007. a01_045). [62] Halmshaw, R. (1954). "The use and scope of Iridium 192 for the radiography of steel". British Journal of Applied Physics 5: 238–243. doi: 10.1088/0508-3443/5/7/302 (http:/ / dx. doi. org/ 10. 1088/ 0508-3443/ 5/ 7/ 302). [63] Hellier, Chuck (2001). Handbook of Nondestructive Evlaluation. The McGraw-Hill Companies. [64] The definition of the meter was changed again in 1983. The meter is currently defined as the distance traveled by light in a vacuum during a time interval of 1⁄299,792,458 of a second. [65] Penzes, W. B. (2001). " Time Line for the Definition of the Meter (http:/ / www. mel. nist. gov/ div821/ museum/ timeline. htm)". National Institute for Standards and Technology. . Retrieved 2008-09-16. [66] General section citations: Recalibration of the U.S. National Prototype Kilogram, R. S. Davis, Journal of Research of the National Bureau of Standards, 90, No. 4, July–August 1985 ( 5.5 MB PDF, here (http:/ / nvl. nist. gov/ pub/ nistpubs/ jres/ 090/ 4/ V90-4. pdf)); and The Kilogram and Measurements of Mass and Force, Z. J. Jabbour et al., J. Res. Natl. Inst. Stand. Technol. 106, 2001, 25–46 ( 3.5 MB PDF, here (http:/ / nvl. nist. gov/ pub/ nistpubs/ jres/ 106/ 1/ j61jab. pdf)) [67] Ziegler,, E.; Hignette, O.; Morawe, Ch.; Tucoulou, R. (2001). "High-efficiency tunable X-ray focusing optics using mirrors and laterally-graded multilayers". Nuclear Instruments and Methods in Physics Research Section A: Accelerators, Spectrometers, Detectors and Associated Equipment 467–468: 954–957. doi: 10.1016/S0168-9002(01)00533-2 (http:/ / dx. doi. org/ 10. 1016/ S0168-9002(01)00533-2). [68] " Face-to-Face with Jerry Johnston, CXC Program Manager & Bob Hahn, Chief Engineer at Optical Coating Laboratories, Inc., Santa Rosa, CA (http:/ / chandra. harvard. edu/ press/ bios/ johnston. html)". Harvard-Smithsonian Center for Astrophysics; Chandra X-ray Center. 1995. . Retrieved 2008-09-24. [69] Möhl, D. (1997). "Production of low-energy antiprotons". Zeitschrift Hyperfine Interactions 109: 33–41. doi: 10.1023/A:1012680728257 (http:/ / dx. doi. org/ 10. 1023/ A:1012680728257). [70] Janowicz, A. H.; Bergman, R. G. (1982). "Carbon-hydrogen activation in completely saturated hydrocarbons: direct observation of M + R-H -> M(R)(H)". Journal of the American Chemical Society 104 (1): 352–354. doi: 10.1021/ja00365a091 (http:/ / dx. doi. org/ 10. 1021/ ja00365a091).
15
Iridium [71] Hoyano, J. K.; Graham, W. A. G. (1982). "Oxidative addition of the carbon-hydrogen bonds of neopentane and cyclohexane to a photochemically generated iridium(I) complex". Journal of the American Chemical Society 104 (13): 3723–3725. doi: 10.1021/ja00377a032 (http:/ / dx. doi. org/ 10. 1021/ ja00377a032). [72] Källström, K; Munslow, I; Andersson, P. G. (2006). "Ir-catalysed asymmetric hydrogenation: Ligands, substrates and mechanism". Chemistry – A European Journal 12 (12): 3194–3200. doi: 10.1002/chem.200500755 (http:/ / dx. doi. org/ 10. 1002/ chem. 200500755). PMID 16304642. [73] Roseblade, S. J.; Pfaltz, A. (2007). "Iridium-catalyzed asymmetric hydrogenation of olefins". Accounts of Chemical Research 40 (12): 1402–1411. doi: 10.1021/ar700113g (http:/ / dx. doi. org/ 10. 1021/ ar700113g). PMID 17672517. [74] Wang, X.; Andersson, M. R.; Thompson, M. E.; Inganäsa, O. (2004). "Electrophosphorescence from substituted poly(thiophene) doped with iridium or platinum complex". Thin Solid Films 468 (1–2): 226–233. doi: 10.1016/j.tsf.2004.05.095 (http:/ / dx. doi. org/ 10. 1016/ j. tsf. 2004. 05. 095). [75] Tonzetich, Zachary J. (2002). " Organic Light Emitting Diodes—Developing Chemicals to Light the Future (http:/ / sa. rochester. edu/ jur/ issues/ fall2002/ tonzetich. pdf)" (PDF). Journal of Undergraduate Research (Rochester University) 1 (1). . Retrieved 2008-10-10. [76] Holder, E.; Langefeld, B. M. W.; Schubert, U. S. (2005-04-25). "New Trends in the Use of Transition Metal-Ligand Complexes for Applications in Electroluminescent Devices". Advanced Materials 17 (9): 1109–1121. doi: 10.1002/adma.200400284 (http:/ / dx. doi. org/ 10. 1002/ adma. 200400284). [77] Mottishaw, J. (1999). " Notes from the Nib Works—Where's the Iridium? (http:/ / www. pencollectors. com/ nib-corner. html)". The PENnant XIII (2). . [78] Crookes, W., ed (1867). "The Paris Exhibition". The Chemical News and Journal of Physical Science XV: 182. [79] Pepper, J. H. (1861). The Playbook of Metals: Including Personal Narratives of Visits to Coal, Lead, Copper, and Tin Mines, with a Large Number of Interesting Experiments Relating to Alchemy and the Chemistry of the Fifty Metallic Elements. Routledge, Warne, and Routledge. p. 455. [80] " Radioisotope Brief: Iridium-192 (Ir-192) (http:/ / emergency. cdc. gov/ radiation/ isotopes/ pdf/ iridium. pdf)" (PDF). Radiation Emergencies. Center for Disease Control and Prevention. 2004-08-18. . Retrieved 2008-09-20.
External links pnb:میڈیرا
16
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17
Platinum
1
Platinum iridium ← platinum → gold Pd ↑ Pt ↓ Ds Periodic Table Extended Periodic Table General Name, symbol, number
platinum, Pt, 78
Element category
transition metals
Group, Period, Block
10, 6, d
Appearance
grayish white
Standard atomic weight
−1
195.084(9) g·mol
Electron configuration
14
[Xe] 4f
Electrons per shell
9
1
5d 6s
2, 8, 18, 32, 17, 1 Physical properties
Phase
solid Density (near r.t.)
21.45 g·cm−3
Liquid density at m.p.
19.77 g·cm−3
Melting point
2041.4 K (1768.3 °C, 3214.9 °F) Boiling point
4098 K (3825 °C, 6917 °F)
Heat of fusion
22.17 kJ·mol−1
Heat of vaporization
469 kJ·mol−1
Specific heat capacity
(25 °C) 25.86 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
2330
(2550)
2815
3143
3556
4094
Platinum
2 Atomic properties Crystal structure
Oxidation states
face centered cubic 6, 5, 4, 3 , 2, 1, -1, -2 (mildly basic oxide)
Electronegativity Ionization energies
2.28 (Pauling scale) 1st: 870 kJ/mol 2nd: 1791 kJ/mol
Atomic radius Covalent radius Van der Waals radius
139 pm 136±5 pm 175 pm Miscellaneous
Magnetic ordering
paramagnetic
Electrical resistivity
(20 °C) 105 n Ω·m
Thermal conductivity
(300 K) 71.6 W·m
Thermal expansion
(25 °C) 8.8 µm·m
Speed of sound (thin rod)
(r.t.) 2800 m·s
Tensile strength
125-240 MPa
Young's modulus
168 GPa
Shear modulus
61 GPa Bulk modulus
Poisson ratio
−1
230 GPa 0.38
Mohs hardness
4–4.5
Vickers hardness
549 MPa
Brinell hardness CAS registry number
392 MPa 7440-06-4 Most-stable isotopes
−1
−1
−1
−1
·K
·K
Platinum
3
Main article: Isotopes of platinum iso 190
Pt
191
Pt
192
Pt
193
Pt
193m
Pt
NA 0.014% syn 0.782% syn syn
half-life 6.5×1011 y 2.76 d
50 y
195
33.832%
195
Pt
196
Pt
197
Pt
197m
Pt
198
Pt
syn 25.242%
α
3.18
186
ε
?
191
?
193
ε
4.33 d 194
195m
DP Os Ir
Pt is stable with 114 neutron
32.967%
Pt
DE (MeV)
192
194
Pt
DM
IT
Ir
0.1355e
193
0.1297e
195
Pt
Pt is stable with 116 neutron Pt is stable with 117 neutron
4.02 d
IT
Pt
196
Pt is stable with 118 neutron
syn
19.8913 h
β
0.719
197
syn
1.59 h
IT
0.3465
197
7.163%
−
Au Pt
198
Pt is stable with 120 neutron References
Platinum (pronounced /ˈplætɨnəm/) is a chemical element with the chemical symbol Pt and an atomic number of 78. Its name is derived from the Spanish term platina del Pinto, which is literally translated into "little silver of the Pinto River."[1] It is in Group 10 of the periodic table of elements. A dense, malleable, ductile, precious, gray-white transition metal, platinum is resistant to corrosion and occurs in some nickel and copper ores along with some native deposits. Platinum is used in jewelry, laboratory equipment, electrical contacts and electrodes, platinum resistance thermometers, dentistry equipment, and catalytic converters. Platinum bullion has the ISO currency code of XPT. Platinum is a commodity with a value that fluctuates according to market forces. On June 5, 2009, Platinum was worth $1263.00 per troy ounce (approximately $40.09 per gram).[2]
Characteristics As a pure metal, platinum is silvery-white in appearance, lustrous, ductile, and malleable.[3] It does not oxidize at any temperature, although it is corroded by halogens, cyanides, sulfur, and caustic alkalis. Platinum is insoluble in hydrochloric and nitric acid, but dissolves in aqua regia to form chloroplatinic acid, H2PtCl6.[4] Platinum's wear- and tarnish-resistance characteristics are well suited for making fine jewelry. Platinum is more precious than gold or silver. Platinum possesses high resistance to chemical attack, excellent high-temperature characteristics, and stable electrical properties. All of these properties have been exploited for industrial applications.
Platinum
4
Isotopes Platinum has six naturally occurring isotopes: 190Pt, 192Pt, 194Pt, 195Pt, 196Pt, and 198Pt. The most abundant of these is 195Pt, comprising 33.83% of all platinum. 190Pt is the least abundant at only .01%. Of the naturally occurring isotopes, only 190Pt is unstable, though it decays with a half-life of 650 × 109 years. 198Pt undergoes alpha decay, but because its half-life is estimated as being greater than 320 × 1012 years, it is considered stable. Platinum also has 31 synthetic isotopes ranging in atomic mass from 166 to 202, making the total number of known isotopes 37. The least stable of these is 166Pt with a half-life of 300 µs, while the most stable is 193Pt with a half-life of 50 years. Most of platinum's isotopes decay by some combination of beta decay and alpha decay. 188Pt, 191Pt, and 193Pt decay primarily by electron capture. 190Pt and 198Pt have double beta decay paths.[5]
Chemistry and compounds Platinum's most common oxidation states are +2, and +4. The +1 and +3 oxidation states are less common, and are often stabilized by metal bonding in bimetallic (or polymetallic) species. As is expected, tetracoordinate platinum(II) compounds tend to adopt a square planar geometry. While elemental platinum is generally unreactive, it dissolves in aqua regia to give soluble hexachloroplatinic acid ("H2PtCl6", formally (H3O)2PtCl6·nH2O ):[6] Pt + 4 HNO3 + 6 HCl → H2PtCl6 + 4 NO2 + 4 H2O This compound has various applications in photography, zinc etchings, indelible ink, plating, mirrors, porcelain coloring, and as a catalyst.[7] Treatment of hexachloroplatinic acid with an ammonium salt, such as ammonium chloride, gives ammonium hexachloroplatinate,[8] which is very insoluble in ammonium solutions. Heating the ammonium salt in the presence of hydrogen reduces it to elemental platinum. Platinum is often isolated from ores and recycled thus.[9] Potassium hexachloroplatinate is similarly insoluble, such that the acid has been used in the determination of potassium ions by gravimetry.[10] When hexachloroplatinic acid is heated, it decomposes through platinum(IV) chloride and platinum(II) chloride to elemental platinum, although the reactions do not occur stepwise, cleanly:[11] (H3O)2PtCl6·n H2O PtCl4
PtCl2 + Cl2
PtCl2
Pt + Cl2
PtCl4 + 2 HCl + (n + 2) H2O
All three reactions are reversible. Platinum(II) and platinum(IV) bromides are known as well. Platinum hexafluoride is a strong oxidizer. Platinum(IV) oxide, PtO2, also known as Adams' Catalyst, is a black powder which is soluble in KOH solutions and concentrated acids.[12] PtO2 and the less common PtO both decompose upon heating.[3] Platinum(II,IV) oxide, Pt3O4, is formed in the following reaction: 2 Pt2+ + Pt4+ + 4 O2− → Pt3O4 Platinum also forms a trioxide, which is actually in the +4 oxidation state. Unlike palladium acetate, platinum(II) acetate is not commercially available. Where a base is desired, the halides have been used in conjunction with sodium acetate.[13] The use of platinum(II) acetylacetonate has also been reported.[14]
Platinum
5
Zeise's salt, containing an ethylene ligand, was one of the first organometallic compounds discovered. Dichloro(cycloocta-1,5-diene)platinum(II) is a commercially available olefin complex, which contains easily displaceable cod ligands ("cod" being an abbreviation of 1,5-cyclooctadiene). The cod complex and the halides are convenient starting points to platinum chemistry. As a soft acid, platinum has a great affinity for sulfur, such as on DMSO; numerous DMSO complexes have been reported and care should be taken in the choice of reaction solvent.[13] Cisplatin, or cis-diamminedichloroplatinum(II) is the first of a series of square planar platinum(II)-containing chemotherapy drugs, including carboplatin and oxaliplatin. These compounds are capable of crosslinking DNA and kill cells by similar pathways to alkylating chemotherapeutic agents.[15]
The hexachloroplatinate ion
Cisplatin
Dichloro(cycloocta-1,5-diene)platinum(II)
Several barium platinides have been synthesized, in which platinum exhibits negative oxidation states ranging from −1 to −2. These include BaPt, Ba3Pt2, and Ba2Pt.[16] Caesium platinide, Cs2Pt, has been shown to contain Pt2− anions.[17] Platinum is also shown to exhibit negative oxidation states at surfaces reduced electrochemically.[18] The negative oxidation states exhibited by platinum, which are unusual for metallic elements, are believed to be due to the relativistic stabilization of the 6s orbitals.[17]
Occurrence Platinum is an extremely rare metal, occurring as only 0.003 ppb in the Earth's crust. It is sometimes mistaken for silver (Ag) but platinum is whiter in appearance. Platinum is often found chemically uncombined as native platinum and alloyed with iridium as platiniridium. Most often the native platinum is found in secondary deposits, platinum is combined with the other platinum group metals in alluvial deposits. The alluvial deposits used by pre-Columbian people in the Platinum ore, US cent included for Chocó Department, Colombia are still a source for scale platinum group metals. Another large alluvial deposit was found in the Ural mountains, Russia, which is still mined.
Platinum
In nickel and copper deposits platinum group metals occur as sulfides (i.e. (Pt,Pd)S)), tellurides (i.e. PtBiTe), antimonides (PdSb), and arsenides (i.e. PtAs2), and as end alloys with nickel or copper. Platinum arsenide, sperrylite (PtAs2), is a major source of platinum associated with nickel ores in the Sudbury Basin Platinum output in 2005 deposit in Ontario, Canada. The rare sulfide mineral cooperite, (Pt,Pd,Ni)S, contains platinum along with palladium and nickel. Cooperite occurs in the Merensky Reef within the Bushveld complex, Gauteng, South Africa.[19] The largest known primary reserves are in the Bushveld complex in South Africa.[20] The large copper–nickel deposits near Norilsk in Russia, and the Sudbury Basin, Canada, are the two other large deposits. In the Sudbury Basin the huge quantities of nickel ore processed makes up for the fact that platinum is present as only 0.5 ppm in the ore. Smaller reserves can be found in the United States,[20] for example in the Absaroka Range in Montana.[21] This is also shown in the production of 2005. In 2005, South Africa was the top producer of platinum with an almost 80% share followed by Russia and Canada.[22] Platinum exists in higher abundances on the Moon and in meteorites. Correspondingly, platinum is found in slightly higher abundances at sites of bolide impact on the Earth that are associated with resulting post-impact volcanism, and can be mined economically; the Sudbury Basin is one such example.
Production Platinum together with the rest of the platinum metals is obtained commercially as a by-product from nickel and copper mining and processing. During electrorefining of copper, noble metals such as silver, gold and the platinum group metals as well as selenium and tellurium settle to the bottom of the cell as anode mud, which forms the starting point for the extraction of the platinum group metals.[23] [24]
If pure platinum is found in placer deposits 1000 cubic centimeters of 99.9% pure platinum or other ores it is isolated from them by various methods of subtracting impurities. Because platinum is significantly denser than many of its impurities, the lighter impurities can be removed by simply floating them away in a water bath. Platinum is also non-magnetic, while nickel and iron are both magnetic. These two impurities are thus removed by running an electromagnet over the mixture. Because platinum has a higher melting point than most other substances, many impurities can be burned or melted away without melting the platinum. Finally, platinum is resistant to hydrochloric and sulfuric acids, while other substances are readily attacked by them. Metal impurities can be removed by stirring the mixture in either of the two acids and recovering the remaining platinum.[25]
6
Platinum
7
One suitable method for purification for the raw platinum, which contains platinum, gold, and the other platinum group metals, is to process it with aqua regia, in which palladium, gold and platinum are dissolved, while osmium, iridium, ruthenium and rhodium stay unreacted. The gold is precipitated by the addition of iron(III) chloride and after filtering of the gold, the platinum is precipitated by the addition of ammonium chloride as ammonium chloroplatinate. Ammonium chloroplatinate can be converted to the metal by heating.[26]
Applications
Cross section of a Metal-core Converter
Of the 239 tonnes of platinum sold in 2006, 130 tonnes were used for automobile emissions control devices, 49 tonnes were used for jewelry, 13.3 tonnes were used in electronics, and 11.2 tonnes were used by the chemical industry as a catalyst. The remaining 35.5 tonnes produced were used in various other minor applications, such as electrodes, anticancer drugs, oxygen sensors, spark plugs and turbine engines.[27]
Catalysis The most common use of platinum is as catalyst in chemical reactions. It has been employed in this application since the early 1800s, when platinum powder was used to catalyze the ignition of hydrogen. The most important application of platinum is in automobiles as a catalytic converter, which allows the complete combustion of low concentrations of unburned hydrocarbon from the exhaust into carbon dioxide and water vapor. Platinum is also used in the petroleum industry as a catalyst in a number of separate processes, but especially in catalytic reforming of straight run naphthas into higher octane gasoline which becomes rich in aromatic compounds. PtO2, also known as Adams' catalyst, is used as a hydrogenation catalyst, specifically for vegetable oils.[7] Platinum metal also strongly catalyzes the decomposition of hydrogen peroxide into water and oxygen gas.[28]
Standard From 1889 to 1960, the meter was defined as the length of a platinum-iridium (90:10) alloy bar, known as the International Prototype Meter bar. The previous bar was made of platinum in 1799. The International Prototype Kilogram remains defined by a cylinder of the same platinum-iridium alloy made in 1879. The standard hydrogen electrode also utilizes a platinized platinum electrode due to its corrosion resistance, and other attributes.
International Prototype Meter bar
Platinum
8
Precious metal Platinum is a precious metal commodity; its bullion has the ISO currency code of XPT. Coins, bars, and ingots are traded or collected. Platinum finds use in jewelry, usually as a 90-95% alloy, due to its inertness and shine. In watchmaking, Vacheron Constantin, Patek Philippe, Rolex, Breitling and other companies use platinum for producing their limited edition watch series. Watchmakers highly appreciate the unique properties of platinum as it neither tarnishes nor wears out.[29]
Platinum Eagle
Price The price of platinum, like other industrial commodities, is more volatile than that of gold. In 2008 the price of platinum ranged from $774 to $2,252 per oz.[31] During
periods
of
sustained
economic
stability and growth, the price of platinum tends to be as much as twice the price of gold, whereas during periods of economic uncertainty[32] , the price of platinum tends Average price of platinum from 1991 to 2007 in US$ [30] per troy ounce (~$40/g). to decrease due to reduced industrial demand, falling below the price of gold. Gold prices are more stable in slow economic times, as gold is considered a safe haven and gold demand is not driven by industrial uses. In the 18th century, platinum's rarity made King Louis XV of France declare it the only metal fit for a king.[33]
Other uses In the laboratory, platinum wire is used for electrodes; platinum pans are used in thermogravimetric analysis. Platinum is used as an alloying agent for various metal products, including fine wires, noncorrosive laboratory containers, medical instruments, dental prostheses, electrical contacts, and thermocouples. Platinum-cobalt, an alloy comprised of roughly 3 parts platinum and 1 part cobalt, is used to make extremely strong permanent magnets.[7] Platinum-based anodes are used in ships, pipelines, and steel piers.[4]
Platinum
9
Symbol of prestige Platinum's rarity as a metal has caused advertisers to associate it with exclusivity and wealth. "Platinum" debit cards have greater privileges than do "gold" ones. "Platinum awards" are the second highest possible, ranking above "gold", "silver" and "bronze", but below Diamond. For example, in the United States a musical album that has sold more than 1,000,000 copies, will be An assortment of native platinum credited as "platinum", whereas an album that sold nuggets more than 10,000,000 copies will be certified as “diamond”. Some products, such as blenders and vehicles, with a silvery-white color are identified as "platinum". Platinum is considered a precious metal, although its use is not as common as the use of gold or silver. The frame of the Crown of Queen Elizabeth the Queen Mother, manufactured for her Coronation as Consort of King George VI, is made of platinum. It was the first British crown to be made of this particular metal.
History Platinum occurs naturally in the alluvial sands of various rivers, though there is little evidence of its use by ancient peoples. However, the metal was used by pre-Columbian Americans near modern-day Esmeraldas, Ecuador to produce artifacts of a white gold-platinum alloy. The first European reference to platinum appears in 1557 in the writings of the Italian humanist Julius Caesar Scaliger as a description of an unknown noble metal found between Darién and Mexico, "which no fire nor any Spanish artifice has yet been able to liquefy."[34]
The alchemical symbol for platinum (shown above) was made by joining the symbols of silver and gold.
In 1741, Charles Wood, a British metallurgist, found various samples of Columbian platinum in Jamaica, which he sent to William Brownrigg for further investigation. Antonio de Ulloa, also credited with the discovery of platinum, returned to Spain from the French Geodesic Mission in 1746 after having been there for eight years. His historical account of the expedition included a description of platinum as being neither separable nor calcinable. Ulloa also anticipated the discovery of platinum mines. After publishing the report in 1748, Ulloa did not continue to investigate the new metal. In 1758, he was sent to superintend mercury mining operations in Huancavelica.[34]
In 1750, after studying the platinum sent to him by Wood, Brownrigg presented a detailed account of the metal to the Royal Society, mentioning that he had seen no mention of it in any previous accounts of known minerals. Brownrigg also made note of platinum's extremely high melting point and refractoriness toward borax. Other chemists across Europe soon began studying platinum, including Torbern Bergman, Jöns Jakob Berzelius, William Lewis, and Pierre Macquer. In 1752, Henrik Scheffer published a detailed scientific description of the metal, which he referred to as "white gold", including an account of how he succeeded in fusing platinum ore with the aid of arsenic. Scheffer described platinum as being less pliable than gold, but with similar resistance to corrosion.[34]
Platinum Carl von Sickingen researched platinum extensively in 1772. He succeeded in making malleable platinum by alloying it with gold, dissolving the alloy in aqua regia, precipitating the platinum with ammonium chloride, igniting the ammonium chloroplatinate, and hammering the resulting finely divided platinum to make it cohere. Franz Karl Achard made the first platinum crucible in 1784. He worked with the platinum by fusing it with arsenic, then later volatilizing the arsenic.[34] In 1786, Charles III of Spain provided a library and laboratory to Pierre-François Chabaneau to aid in his research of platinum. Chabaneau succeeded in removing various impurities from the ore, including gold, mercury, lead, copper, and iron. This led him to believe that he was working with a single metal, but in truth the ore still contained the yet-undiscovered platinum group metals. This led to inconsistent results in his experiments. At times the platinum seemed malleable, but when it was alloyed with iridium, it would be much more brittle. Sometimes the metal was entirely incombustible, but when alloyed with osmium, it would volatilize. After several months, Chabaneau succeeded in producing 23 kilograms of pure, malleable platinum by hammering and compressing the sponge form while white-hot. Chabeneau realized that the infusibility of platinum would lend value to objects made of it, and so started a business with Joaquín Cabezas producing platinum ingots and utensils. This started what is known as the "platinum age" in Spain.[34] From 1875 to 1960 the SI unit of length (the standard meter) was defined as the distance between two lines on a standard bar of an alloy of ninety percent platinum and ten percent iridium, measured at 0 degrees Celsius. In 2007 Gerhard Ertl won the Nobel Prize in Chemistry for determining the detailed molecular mechanisms of the catalytic oxidation of carbon monoxide over platinum (catalytic converter).
Precautions According to the Centers for Disease Control and Prevention, short-term exposure to platinum salts "may cause irritation of the eyes, nose, and throat" and long-term exposure "may cause both respiratory and skin allergies." The current OSHA standard is 0.002 milligram per cubic meter of air averaged over an 8-hour work shift.[35] Certain platinum complexes are used in chemotherapy and show good anti-tumor activity for some tumors. Cisplatin is particularly effective against testicular cancer; cure rate was improved from 10% to 85%.[36] However, the side effects are severe. Cisplatin causes cumulative, irreversible kidney damage and deafness.[37] As with other ototoxic agents, deafness may be secondary to interactions with melanin in the stria vascularis. As platinum is a catalyst in the manufacture of the silicone rubber and gel components of several types of medical implants (breast implants, joint replacement prosthetics, artificial lumbar discs, vascular access ports), the possibility that platinum free radicals could enter the body and cause adverse effects has merited study. The FDA and other countries have reviewed the issue and found no evidence to suggest toxicity in vivo.[38]
10
Platinum
See also • • • • • • • •
Platinum black Platinum coin Platinum group Platinum in Africa Merensky Reef Precious metal Palladium Platinum nanoparticles
References • Nuclides and Isotopes [39] Fourteenth Edition: Chart of the Nuclides, General Electric Company, 1989. • Jefferson Lab — The Element Platinum [40]
External links • • • • • • •
The Platinum Group Metals Database [41] A balanced historical account of the sequence of discoveries of platinum; illustrated. [42] WebElements.com — Platinum [43] Platinum Metals Review E-Journal [44] Platinum Guild International [45] United States Geological Survey Platinum-Group Metals Statistics and Information [46] picture of a 999.5 fine platinum ingot in the element collection from Heinrich Pniok [47]
pnb:منیٹالپ
References [1] Woods, Ian (2004). The Elements: Platinum. Benchmark Books. ISBN 978-0761415503. [2] " Live Market Quotes (http:/ / www. kitco. com/ market/ )". Kitco. . Retrieved 2009-04-03. [3] Lagowski, J. J., ed (2004). Chemistry Foundations and Applications. 3. Thomson Gale. pp. 267–268. ISBN 0-02-865724-1. [4] CRC contributors (2007–2008). "Platinum". in Lide, David R.. CRC Handbook of Chemistry and Physics. 4. New York: CRC Press. pp. 26. ISBN 978-0-8493-0488-0. [5] Audi, G. (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3–128. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [6] George B. Kauffman (1967). "Ammonium Hexachloroplatinate(IV)". Inorganic Syntheses 9: 182–185. doi: 10.1002/9780470132401.ch51 (http:/ / dx. doi. org/ 10. 1002/ 9780470132401. ch51). [7] Krebs, Robert E. (1998). "Platinum". The History and Use of our Earth's Chemical Elements. Greenwood Press. pp. 124–127. ISBN 0-313-30123-9. [8] George B. Kauffman (1967). "Ammonium Hexachloroplatinate(IV)". Inorganic Syntheses 9: 182–185. doi: 10.1002/9780470132401.ch51 (http:/ / dx. doi. org/ 10. 1002/ 9780470132401. ch51). [9] Cotton, S. A. Chemistry of Precious Metals, Chapman and Hall (London): 1997. ISBN 0-7514-0413-6. [10] G. F. Smith, J. L. Gring (1933). "The Separation and Determination of the Alkali Metals Using Perchloric Acid. V. Perchloric Acid and Chloroplatinic Acid in the Determination of Small Amounts of Potassium in the Presence of Large Amounts of Sodium". J. Am. Chem. Soc. 55 (10): 3957–3961. doi: 10.1021/ja01337a007 (http:/ / dx. doi. org/ 10. 1021/ ja01337a007). [11] A. E. Schweizer, G. T. Kerr (1978). "Thermal Decomposition of Hexachloroplatinic Acid". Inorg. Chem. 17 (8): 2326–2327. doi: 10.1021/ic50186a067 (http:/ / dx. doi. org/ 10. 1021/ ic50186a067). [12] Perry, D. L. (1995). Handbook of Inorganic Compounds. CRC Press. pp. 296–298. ISBN 0-8492-8671-3.
11
Platinum [13] Y. Han, H. V. Huynh, G. K. Tan (2007). "Mono- vs Bis(carbene) Complexes: A Detailed Study on Platinum(II)−Benzimidazolin-2-ylidenes". Organometallics 26: 4612. doi: 10.1021/om700543p (http:/ / dx. doi. org/ 10. 1021/ om700543p). [14] Sebastian Ahrens and Thomas Strassner (2006). "Detour-free synthesis of platinum-bis-NHC chloride complexes, their structure and catalytic activity in the CH activation of methane". Inorganica Chimica Acta 359: 4789. doi: 10.1016/j.ica.2006.05.042 (http:/ / dx. doi. org/ 10. 1016/ j. ica. 2006. 05. 042). [15] Richards, A.D.; Rodger, A. (2007). " Synthetic metallomolecules as agents for the control of DNA structure (http:/ / www. rsc. org/ publishing/ journals/ CS/ article. asp?doi=b609495c)". Chem. Soc. Rev. 36: 471–483. doi: 10.1039/b609495c (http:/ / dx. doi. org/ 10. 1039/ b609495c). . [16] Karpov, Andrey; Konuma, M; Jansen, M (Feb 2006). "An experimental proof for negative oxidation states of platinum: ESCA-measurements on barium platinides". Chemical communications (Cambridge, England) (8): 838–40. doi: 10.1039/b514631c (http:/ / dx. doi. org/ 10. 1039/ b514631c). ISSN 1359-7345 (http:/ / worldcat. org/ issn/ 1359-7345). PMID 16479284. edit (http:/ / en. wikipedia. org/ wiki/ Template:cite_doi/ 10. 1039. 2fb514631c) [17] Jansen, Martin (June 2005). "Effects of relativistic motion of electrons on the chemistry of gold and platinum". Solid State Sciences 7 (12): 1464–1474. doi: 10.1016/j.solidstatesciences.2005.06.015 (http:/ / dx. doi. org/ 10. 1016/ j. solidstatesciences. 2005. 06. 015). edit (http:/ / en. wikipedia. org/ wiki/ Template:cite_doi/ 10. 1016. 2fj. solidstatesciences. 2005. 06. 015) [18] Ghilane, J.; C. Lagrost; M. Guilloux-Viry; J. Simonet; M. Delamar; C. Mangeney; P. Hapiot (March 2007). "Spectroscopic Evidence of Platinum Negative Oxidation States at Electrochemically Reduced Surfaces". The Journal of Physical Chemistry C 111 (15): 5701–5707. doi: 10.1021/jp068879d (http:/ / dx. doi. org/ 10. 1021/ jp068879d). edit (http:/ / en. wikipedia. org/ wiki/ Template:cite_doi/ 10. 1021. 2fjp068879d) [19] Xiao, Z.; Laplante, A. R. (2004). "Characterizing and recovering the platinum group minerals—a review". Minerals Engineering 17: 961–979. doi: 10.1016/j.mineng.2004.04.001 (http:/ / dx. doi. org/ 10. 1016/ j. mineng. 2004. 04. 001). [20] Seymour, R. J.; O'Farrelly, J. I. (2001). "Platinum-group metals". Kirk Othmer Encyclopedia of Chemical Technology. Wiley. doi: 10.1002/0471238961.1612012019052513.a01.pub2 (http:/ / dx. doi. org/ 10. 1002/ 0471238961. 1612012019052513. a01. pub2). [21] " Mining Platinum in Montana (http:/ / query. nytimes. com/ gst/ fullpage. html?res=9802E3D6153AF930A2575BC0A96E958260)". New York Times. 1998-08-13. . Retrieved 2008-09-09. [22] " Platinum–Group Metals (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ platinum/ platimcs07. pdf)" (PDF). U.S. Geological Survey, Mineral Commodity Summaries. January 2007. . Retrieved 2008-09-09. [23] George, M. W. (2008). " Platinum-group metals (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ platinum/ mcs-2008-plati. pdf)" (pdf). U.S. Geological Survey Mineral Commodity Summaries (USGS Mineral Resources Program). . [24] George, M. W.. 2006 Minerals Yearbook: Platinum-Group Metals (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ platinum/ myb1-2006-plati. pdf). United States Geological Survey USGS. . [25] Heiserman, David L. (1992). Exploring Chemical Elements and their Compounds. TAB Books. pp. 272–274. ISBN 0-8306-3018-X. [26] Hunt, L. B.; Lever, F. M. (1969). " Platinum Metals: A Survey of Productive Resources to industrial Uses (http:/ / www. platinummetalsreview. com/ pdf/ pmr-v13-i4-126-138. pdf)". Platinum Metals Review 13 (4): 126–138. . Retrieved 2009-10-02. [27] George, Micheal W.. " Mineral Yearbook 2006: Platinum-Group Metals (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ platinum/ myb1-2006-plati. pdf)". United States Geological Survey. . Retrieved 2008-09-25. [28] Petrucci, Ralph H. (2007). General Chemistry: Principles & Modern Applications (9th ed.). Prentice Hall. pp. 606. ISBN 0131493302. [29] " Unknown Facts about Platinum (http:/ / watches. infoniac. com/ index. php?page=post& id=44)". watches.infoniac.com. . Retrieved 2008-09-09. [30] " London Platinum and Palladium Market (http:/ / www. lppm. org. uk/ statistics_cover. html)". The London Platinum and Palladium Market. . Retrieved 2008-08-08. [31] " One Year Platinum (http:/ / www. kitco. com/ charts/ popup/ pt0365nyb. html)". Kitco. . Retrieved 2009-01-12. [32] " Platinum versus Gold (http:/ / www. speculative-investor. com/ new/ article150402. html)". The Speculative Invertor. . [33] " Platinum (http:/ / www. mineralszone. com/ minerals/ platinum. html)". Minerals Zone. . Retrieved 2008-09-09. [34] Weeks, M. E. (1968). Discovery of the Elements (7 ed.). Journal of Chemical Education. pp. 385–407. ISBN 0848685792. OCLC 23991202 (http:/ / worldcat. org/ oclc/ 23991202).
12
Platinum [35] " Occupational Health Guideline for Soluble Platinum Salts (as Platinum) (http:/ / www. cdc. gov/ niosh/ pdfs/ 0520. pdf)" (PDF). Centers for Disease Control and Prevention. . Retrieved 2008-09-09. [36] Einhorn LH. (01 Nov 1990). " Treatment of testicular cancer: a new and improved model (http:/ / jco. ascopubs. org/ cgi/ content/ abstract/ 8/ 11/ 1777)". J. Clin. Oncol. 8 (11): 1777–81. PMID 1700077. . [37] Von Hoff DD, et al. (1979). " Toxic effects of cis-dichlorodiammineplatinum(II) in man (http:/ / www. ncbi. nlm. nih. gov/ entrez/ query. fcgi?cmd=Retrieve& db=PubMed& list_uids=387223& dopt=Citation)". Cancer Treat. Rep. 63 (9–10): 1527–31. . [38] " FDA Backgrounder on Platinum in Silicone Breast Implants (http:/ / www. fda. gov/ cdrh/ breastimplants/ platinum. html)". U.S. Food and Drug Administration. . Retrieved 2008-09-09. [39] [40] [41] [42] [43] [44] [45] [46] [47]
http:/ / chartofthenuclides. com/ default. html http:/ / education. jlab. org/ itselemental/ ele078. html http:/ / www. platinummetalsreview. com/ jmpgm/ index. jsp http:/ / www. vanderkrogt. net/ elements/ elem/ pt. html http:/ / www. webelements. com/ webelements/ elements/ text/ Pt/ index. html http:/ / www. platinummetalsreview. com/ http:/ / www. platinumguild. com/ http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ platinum/ http:/ / www. pse-mendelejew. de/ bilder/ pt. jpg
13
Article Sources and Contributors
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14
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License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
15
Gold
1
Gold platinum ← gold → mercury Ag ↑ Au ↓ Rg Periodic Table Extended Periodic Table General Name, symbol, number
gold, Au, 79
Element category
transition metals
Group, Period, Block
11, 6, d
Appearance
metallic yellow
Standard atomic weight
−1
196.966569(4) g·mol
Electron configuration
14
[Xe] 4f
Electrons per shell
10
5d
1
6s
2, 8, 18, 32, 18, 1 Physical properties
Phase
solid Density (near r.t.)
19.3 g·cm−3
Liquid density at m.p.
17.31 g·cm−3
Melting point
1337.33 K (1064.18 °C, 1947.52 °F) Boiling point
3129 K (2856 °C, 5173 °F)
Heat of fusion
12.55 kJ·mol−1
Heat of vaporization
324 kJ·mol−1
Specific heat capacity
(25 °C) 25.418 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1646
1814
2021
2281
2620
3078
Atomic properties
Gold
2 Crystal structure
face centered cubic
Oxidation states
-1, 1, 2, 3, 4, 5 (amphoteric oxide)
Electronegativity
2.54 (Pauling scale)
Ionization energies
1st: 890.1 kJ/mol 2nd: 1980 kJ/mol
Atomic radius
144 pm
Covalent radius
136±6 pm
Van der Waals radius
166 pm Miscellaneous
Magnetic ordering
diamagnetic
Electrical resistivity
(20 °C) 22.14 n Ω·m
Thermal conductivity
(300 K) 318 W·m
Thermal expansion
(25 °C) 14.2 µm·m
Speed of sound (thin rod)
(r.t.) (hard-drawn) −1 2030 m·s
Young's modulus
78 GPa
Tensile strain
0.00157
Shear modulus
27 GPa
−1
Bulk modulus
−1
·K
−1
·K
−1
180 GPa
Poisson ratio
0.44
Mohs hardness
2.5
Vickers hardness
216 MPa
Brinell hardness
? 2450 MPa
CAS registry number
7440-57-5 Most-stable isotopes
Main article: Isotopes of gold iso
NA
half-life
DM
DE (MeV)
DP
195
syn
186.10 d
ε
0.227
195
196
syn
6.183 d
ε
1.506
196
β−
0.686
196
Au Au
197
Au
100%
Pt Pt Hg
197
Au is stable with 118 neutron
198
syn
2.69517 d
β−
1.372
198
199
syn
3.169 d
β−
0.453
199
Au Au
References
Hg Hg
Gold
3
Gold (pronounced /ˈɡoʊld/) is a chemical element with the symbol Au (Latin: aurum) and an atomic number of 79. It has been a highly sought-after precious metal in jewelry, in sculpture, and for ornamentation since the beginning of recorded history. The metal occurs as nuggets or grains in rocks, in veins and in alluvial deposits. Gold is dense, soft, shiny and the most malleable and ductile pure metal known. Pure gold has a bright yellow color and luster traditionally considered attractive, which it maintains without oxidizing in air or water. It is one of the coinage metals and formed the basis for the gold standard used before the collapse of the Bretton Woods system in 1971. At the end of 2006, it was estimated that all the gold ever mined totaled 158,000 tonnes. [1] This can be represented by a cube with an edge length of just 20.2 meters. Modern industrial uses include dentistry and electronics, where gold has traditionally found use because of its good resistance to oxidative corrosion and excellent quality as a conductor of electricity. Chemically, gold is a transition metal and can form trivalent and univalent cations upon solvation. At STP it is attacked by aqua regia (a mixture of acids), forming chloroauric acid and by alkaline solutions of cyanide but not by single acids such as hydrochloric, nitric or sulfuric acids. Gold dissolves in mercury, forming amalgam alloys, but does not react with it. Since gold is insoluble in nitric acid which will dissolve silver and base metals, this is exploited as the basis of the gold refining technique known as "inquartation and parting". Nitric acid has long been used to confirm the presence of gold in items, and this is the origin of the colloquial term "acid test", referring to a gold standard test for genuine value.
Characteristics Gold is the most malleable and ductile of all metals; a single gram can be beaten into a sheet of one square meter, or an ounce into 300 square feet. Gold leaf can be beaten thin enough to become translucent. The transmitted light appears greenish blue, because gold strongly reflects yellow and red.[2] Gold readily creates alloys with many other metals. These alloys can be produced to modify the hardness and other metallurgical properties, to control melting point or to create exotic colors (see below). Gold is a good conductor of heat and electricity and reflects infra red radiation strongly. Chemically, it is unaffected by air, moisture and most corrosive reagents, and is therefore well-suited for use in coins and jewelry and as a protective coating on other, more reactive, metals. However, it is not chemically inert. Free halogens will native Gold nuggets react with gold, and aqua regia dissolves it via formation of chlorine gas which attacks gold to form the chloraurate ion. Gold also dissolves in alkaline solutions of potassium cyanide and in mercury, forming a gold-mercury amalgam.
Gold
4 Common oxidation states of gold include +1 (gold(I) or aurous compounds) and +3 (gold(III) or auric compounds). Gold ions in solution are readily reduced and precipitated out as gold metal by adding any other metal as the reducing agent. The added metal is oxidized and dissolves allowing the gold to be displaced from solution and be recovered as a solid precipitate. High quality pure metallic gold is tasteless; in keeping with its resistance to corrosion (it is metal ions which confer taste to metals).
Gold nuggets found in Arizona
In addition, gold is very dense, a cubic meter weighing 19300 kg. By comparison, the density of lead is 11340 kg/m³, and that of the densest element, osmium,
is 22610 kg/m³.
Color of gold The color of pure gold is metallic yellow. Gold, caesium and copper are the only metallic elements with a natural color other than gray or white. The usual gray color of metals depends on their "electron sea" that is capable of absorbing and re-emitting photons over a wide range of frequencies. Gold reacts differently, depending on subtle relativistic effects that affect the orbitals around gold atoms. [3] [4] Common colored gold alloys such as rose gold can be created by the addition of various amounts of copper and silver, as indicated in the Different colors of Ag-Au-Cu alloys diagram above. Alloys containing palladium or nickel are also important in commercial jewelry as these produce white gold alloys. Less commonly, addition of manganese, aluminium, iron, indium and other elements can produce more unusual colors of gold for various applications.[5]
Gold
5
Applications As the metal Medium of monetary exchange In various countries, gold was used as a standard for monetary exchange, but this practice has been abandoned with the rise of fiat currency. The last country to back their money with gold was Switzerland, which backed 40% of its value until it joined the International Monetary Fund in 1999. [6] Pure gold is too soft for ordinary use and is typically hardened by alloying with copper or other base metals. The gold content of gold alloys is measured in carats (k), pure gold being designated as 24k. Gold coins intended for circulation from 1526 into the 1930s were typically a standard 22k alloy called crown gold, for hardness. Modern collector/investment bullion coins (which do not require good mechanical wear properties) are typically 24k, although the American Gold Eagle, the British gold sovereign and the South African Krugerrand continue to be made at 22k, on historical tradition. The special issue Canadian Gold Maple Leaf coin contains the highest purity gold of any bullion coin, at 99.999% (.99999 fine). The popular issue Canadian Gold Maple Leaf coin has a purity of 99.99%. Several other 99.99% pure gold coins are currently available, including Australia's Gold Kangaroos (first appearing in 1986 as the Australian Gold Nugget, with the kangaroo theme appearing in 1989), the several coins of the Australian Lunar Calendar series, and the Austrian Philharmonic. In 2006, the U.S. Mint began production of the American Buffalo gold bullion coin also at 99.99% purity. Special issue Canadian Gold Maple Leaf coin with the highest purity of any gold coin at a guaranteed 99.999%
Gold was used as a medium of monetary exchange throughout history together with or instead of other minerals, like silver, salt, and copper. At the beginning of World War I the warring nations went onto a fractional gold standard, inflating their currencies to finance the war effort. After World War II gold was replaced by a system of convertible currency following the Bretton Woods system. Many holders of gold in storage (as bullion coin or bullion) hold it as a hedge against inflation or other economic disruptions. The ISO currency code of gold bullion is XAU.
Gold Jewelry Because of the softness of pure (24k) gold, it is usually alloyed with base metals for use in jewelry, altering its hardness and ductility, melting point, color and other properties. Alloys with lower caratage, typically 22k, 18k, 14k or 10k, contain higher percentages of copper, or other base metals or silver or palladium in the alloy. Copper is the most commonly used base metal, yielding a redder color. Eighteen carat gold containing 25% copper is found in antique and Russian jewelry and has Moche gold necklace depicting feline heads. Larco Museum Collection. a distinct, though not dominant, copper cast, creating Lima-Peru rose gold. Fourteen carat gold-copper alloy is nearly identical in color to certain bronze alloys, and both may be used to produce police, as well as other, badges. Blue gold can be made by alloying with iron and purple gold can be made by alloying with aluminium, although rarely done except in specialized jewelry. Blue gold is more brittle and therefore more difficult to work with when making jewelry. Fourteen and eighteen carat gold alloys with silver alone appear greenish-yellow and are referred to as green gold. White gold alloys can be made with palladium or nickel. White 18 carat gold containing 17.3% nickel, 5.5% zinc and 2.2% copper is silver in appearance. Nickel is toxic, however, and its release from nickel white gold is controlled by legislation in Europe. Alternative white gold alloys are available based on palladium, silver and other white metals (World Gold Council), but the palladium alloys are more expensive than those using nickel. High-carat white gold alloys are far more resistant to corrosion than are either pure silver or sterling silver. The Japanese craft of Mokume-gane exploits the color contrasts between laminated colored gold alloys to produce decorative wood-grain effects.
Medicine • In medieval times, gold was often seen as beneficial for the health, in the belief that something that rare and beautiful could not be anything but healthy. Even some modern esotericists and forms of alternative medicine assign metallic gold a healing power.[7] Some gold salts do have anti-inflammatory properties and are used as pharmaceuticals in the treatment of arthritis and other similar conditions. [8] However, only salts and radioisotopes of gold are of pharmacological value, as elemental (metallic) gold is inert to all chemicals it encounters inside the body. • In modern times injectable gold has been proven to help to reduce the pain and swelling of rheumatoid arthritis and tuberculosis.[8] [9] • Dentistry. Gold alloys are used in restorative dentistry, especially in tooth restorations, such as crowns and permanent bridges. The gold alloys' slight malleability facilitates the creation of a superior molar mating surface with other teeth and produces results that are generally more satisfactory than those produced by the creation of porcelain crowns. The use of gold crowns in more prominent teeth such as incisors is favored in some cultures and discouraged in others. • Colloidal gold (colloidal sols of gold nanoparticles) in water are intensely red-colored, and can be made with tightly-controlled particle sizes up to a few tens of nm across by reduction of gold chloride with citrate or ascorbate ions. Colloidal gold is used in
6
Gold
7 research applications in medicine, biology and materials science. The technique of immunogold labeling exploits the ability of the gold particles to adsorb protein molecules onto their surfaces. Colloidal gold particles coated with specific antibodies can be used as probes for the presence and position of antigens on the surfaces of cells (Faulk and Taylor 1979). In ultrathin sections of tissues viewed by electron microscopy, the immunogold labels appear as extremely dense round spots at the position of the antigen (Roth et al. 1980). Colloidal gold is also the form of gold used as gold paint on ceramics prior to firing.
• Gold, or alloys of gold and palladium, are applied as conductive coating to biological specimens and other non-conducting materials such as plastics and glass to be viewed in a scanning electron microscope. The coating, which is usually applied by sputtering with an argon plasma, has a triple role in this application. Gold's very high electrical conductivity drains electrical charge to earth, and its very high density provides stopping power for electrons in the SEM's electron beam, helping to limit the depth to which the electron beam penetrates the specimen. This improves definition of the position and topography of the specimen surface and increases the spatial resolution of the image. Gold also produces a high output of secondary electrons when irradiated by an electron beam, and these low-energy electrons are the most commonly-used signal source used in the scanning electron microscope. • The isotope gold-198, (half-life: 2.7 days) is used in some cancer treatments and for treating other diseases.[10]
Food and drink • Gold can be used in food and has the E Number 175.[11] • Gold leaf, flake or dust is used on and in some gourmet foods, notably sweets and drinks as decorative ingredient.[12] Gold flake was used by the nobility in Medieval Europe as a decoration in food and drinks, in the form of leaf, flakes or dust, either to demonstrate the host's wealth or in the belief that something that valuable and rare must be beneficial for one's health. • Goldwasser (English: Goldwater) is a traditional herbal liqueur produced in Gdańsk, Poland, and Schwabach, Germany, and contains flakes of gold leaf. There are also some expensive (~$1000) cocktails which contain flakes of gold leaf[13] . However, since metallic gold is inert to all body chemistry, it adds no taste nor has it any other nutritional effect and leaves the body unaltered.[14]
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Industry • Gold solder is used for joining the components of gold jewelry by high-temperature hard soldering or brazing. If the work is to be of hallmarking quality, gold solder must match the carat weight of the work, and alloy formulas are manufactured in most industry-standard carat weights to color match yellow and white gold. Gold solder is usually made in at least three melting-point ranges referred to as Easy, Medium and Hard. By using the hard, high-melting point solder first, followed by solders with progressively lower melting points, goldsmiths can assemble complex items with several separate soldered joints.
The 220 kg gold brick displayed in Chinkuashi Gold Museum, Taiwan, Republic of China
• Gold can be made into thread and used in embroidery. • Gold is ductile and malleable, meaning it can be drawn into very thin wire and can be beaten into very thin sheets known as gold leaf. • Gold produces a deep, intense red color when used as a coloring agent in cranberry glass. • In photography, gold toners are used to shift the color of silver bromide black and white prints towards brown or blue tones, or to increase their stability. Used on sepia-toned prints, gold toners produce red tones. Kodak published formulas for several types of gold toners, which use gold as the chloride (Kodak, 2006). • As gold is a good reflector of electromagnetic
The world's largest gold bar weighs 250 kg. Toi museum, Japan.
radiation such as infrared and visible light as well as radio waves, it is used for the protective coatings on many artificial satellites, in infrared protective faceplates in thermal protection suits and astronauts' helmets and in electronic warfare planes like the EA-6B Prowler. • Gold is used as the reflective layer on some high-end CDs. • Automobiles may use gold for heat insulation. McLaren uses gold foil in the engine compartment of its F1 model.[15] • Gold can be manufactured so thin that it appears transparent. It is used in some aircraft cockpit windows for de-icing or anti-icing by passing electricity through it. The heat produced by the resistance of the gold is enough to deter ice from forming.[16]
Gold
Electronics • The concentration of free electrons in gold metal is 5.90×1022 cm−3. Gold is highly conductive to electricity, and has been used for electrical wiring in some high energy applications (silver is even more conductive per volume, but gold has the advantage of corrosion resistance). For example, gold electrical wires were used during some of the Manhattan Project's atomic experiments, but large high current silver wires were used in the calutron isotope separator magnets in the project. • Though gold is attacked by free chlorine, its good conductivity and general resistance to oxidation and corrosion in other environments (including resistance to non-chlorinated acids) has led to its widespread A gold nugget of 5 mm in diameter (bottom) can be expanded through industrial use in the electronic era as a thin layer hammering into a gold foil of about 0.5 coating electrical connectors of all kinds, thereby square meter. Toi museum, Japan. ensuring good connection. For example, gold is used in the connectors of the more expensive electronics cables, such as audio, video and USB cables. The benefit of using gold over other connector metals such as tin in these applications is highly debated. Gold connectors are often criticized by audio-visual experts as unnecessary for most consumers and seen as simply a marketing ploy. However, the use of gold in other applications in electronic sliding contacts in highly humid or corrosive atmospheres, and in use for contacts with a very high failure cost (certain computers, communications equipment, spacecraft, jet aircraft engines) remains very common, and is unlikely to be replaced in the near future by any other metal. • Besides sliding electrical contacts, gold is also used in electrical contacts because of its resistance to corrosion, electrical conductivity, ductility and lack of toxicity.[17] Switch contacts are generally subjected to more intense corrosion stress than are sliding contacts. • Fine gold wires are used to connect semiconductor devices to their packages through a process known as wire bonding.
Other • Many competitions, and honors, such as the Olympics and the Nobel Prize, award a gold medal to the winner.
As gold chemical compounds Gold is attacked by and dissolves in alkaline solutions of potassium or sodium cyanide, and gold cyanide is the electrolyte used in commercial electroplating of gold onto base metals and electroforming. Gold chloride (chloroauric acid) solutions are used to make colloidal gold by reduction with citrate or ascorbate ions. Gold chloride and gold oxide are used to make highly-valued cranberry or red-colored glass, which, like colloidal gold sols, contains evenly-sized spherical gold nanoparticles." Colored glass chemistry [18]". http:/ / chemistry. about. com/ cs/ inorganic/ a/ aa032503a. htm. Retrieved 2009-06-06.
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History
The Sun symbol has long represented gold
Gold has been known and highly valued since prehistoric times. It may have been the first metal used by humans and was valued for ornamentation and rituals. Egyptian hieroglyphs from as early as 2600 BC describe gold, which king Tushratta of the Mitanni claimed was "more plentiful than dirt" in Egypt.[19] Egypt and especially Nubia had the resources to make them major gold-producing areas for much of history. The earliest known map is known as the Turin Papyrus Map and shows the plan of a gold mine in Nubia together with indications of the local geology. The primitive working methods are described by Strabo and included fire-setting. Large mines also occurred across the Red Sea in what is now Saudi Arabia. The legend of the golden fleece may refer to the use of
The Turin Papyrus Map
fleeces to trap gold dust from placer deposits in the ancient world. Gold is mentioned frequently in the Old Testament, starting with Genesis 2:11 (at Havilah) and is included with the gifts of the magi in the first chapters of Matthew New Testament. The Book of Revelation 21:21 describes the city of New Jerusalem as having streets "made of pure gold, clear as crystal". The south-east corner of the Black Sea was famed for its gold. Exploitation is said to date from the time of Midas, and this gold was important in the establishment of what is probably the world's earliest coinage in Lydia around 610 BC.[20] From 6th or 5th century BC, Chu (state) circulated Ying Yuan, one kind of square gold coin.
The Romans developed new methods for extracting gold on a large scale using hydraulic mining methods, especially in Spain from 25 BC onwards and in Romania from 150 AD onwards. One of their largest mines was at Las Medulas in León (Spain), where seven long aqueducts enabled them to sluice most of a large alluvial deposit. The mines at Roşia Montană in Transylvania were also very large, and until very recently, still mined by opencast methods. They also exploited smaller deposits in Britain, such as placer and hard-rock deposits at Dolaucothi. The various methods they used are well described by Pliny the Elder in his encyclopedia Naturalis Historia written towards the end of the first century AD. Funerary mask of Tutankhamun
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The Mali Empire in Africa was famed throughout the old world for its large amounts of gold. Mansa Musa, ruler of the empire (1312–1337) became famous throughout the old world for his great hajj to Mecca in 1324. When he passed through Cairo in July 1324, he was reportedly accompanied by a camel train that included thousands of people and nearly a hundred camels. He gave away so much gold that it depressed the price in Egypt for over a decade.[21] A contemporary Arab historian remarked:
Jason returns with the golden fleece on an Apulian red-figure calyx krater, ca. 340–330 BC.
“
Gold was at a high price in Egypt until they came in that year. The mithqal did not go below 25 dirhams and was generally above, but from that time its value fell and it cheapened in price and has remained cheap till now. The mithqal does not exceed 22 dirhams or less. This has been the state of affairs for about twelve years until this day by reason of the large amount of gold which they brought into Egypt and spent there [...]
”
[22]
—Chihab Al-Umari
The European exploration of the Americas was fueled in no small part by reports of the gold ornaments displayed in great profusion by Native American peoples, especially in Central America, Peru, Ecuador and Colombia. Although the price of some platinum group metals can be much higher, gold has long been considered the most desirable of precious metals, and its value has been used as the standard for many currencies (known as the gold standard) in history. Gold has been used as a symbol for purity, value, royalty, and particularly roles that combine these properties. Gold as a sign of wealth and prestige was made fun of by Thomas More in his treatise Utopia. On that imaginary island, gold is so abundant that it is used to make chains for slaves, tableware and lavatory-seats. When ambassadors from other countries arrive, dressed in ostentatious gold jewels and badges, the Utopians mistake them for menial servants, paying homage instead to the most modestly-dressed of their party. There is an age-old tradition of biting gold in order to test its authenticity. Although this is certainly not a professional way of examining gold, the bite test should score the gold because gold is a soft metal, as indicated by its score on the Mohs' scale of mineral hardness. The purer the gold the easier it should be to mark it. Painted lead can cheat this test because lead is softer than gold (and may invite a small risk of lead poisoning if sufficient lead is absorbed by the biting). Gold in antiquity was relatively easy to obtain geologically; however, 75% of all gold ever produced has been extracted since 1910.[23] It has been estimated that all the gold in the world that has ever been refined would form a single cube 20 m (66 ft) on a side (equivalent to 8000 m³).[23]
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One main goal of the alchemists was to produce gold from other substances, such as lead — presumably by the interaction with a mythical substance called the philosopher's stone. Although they never succeeded in this attempt, the alchemists promoted an interest in what can be done with substances, and this laid a foundation for today's chemistry. Their symbol for gold was the circle with a point at its center (☉), which was also the astrological symbol, and the ancient Chinese character, for the Sun. For modern creation of artificial gold by neutron capture, see gold synthesis. During the 19th century, gold rushes occurred whenever large gold deposits were discovered. The first documented discovery of gold in the United States was at the Reed Gold Mine near Georgeville, North Carolina in 1803.[24] The first major gold strike in the United States occurred in a small north Georgia town called Dahlonega.[25] Further gold rushes occurred in California, Colorado, Otago, Australia, Witwatersrand, Black Hills, and Klondike. Because of its historically high value, much of the gold mined throughout history is still in circulation in one form or another.
Occurrence Gold's atomic number of 79 makes it one of the higher
Gold ore
atomic number elements which occur naturally. Like all elements with atomic numbers of six or higher (that is, carbon and elements beyond it), gold is thought to have been formed from a nucleogenesis process beginning from hydrogen in stars which then became supernovas. Their explosions scattered metal-containing dusts (including heavy elements like gold) into the region of space in which they later condensed into our solar system and the Earth. On Earth, whenever elemental gold occurs, it appears most often as a metal solid solution of gold with silver, i.e. a gold silver alloy. Such alloys usually have a silver content of 8–10%. Electrum is elemental gold with more than 20% silver. Electrum's color runs from golden silvery to silvery, dependent upon the silver content. The more silver, the lower the specific gravity.
This 156-ounce (4.85 kg) nugget was found by an individual prospector in the Southern California Desert using a metal detector.
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Relative sizes of a an 860kg rock ore, and the 30g of gold that can be extracted from it. Toi gold mine, Japan.
Gold is found in ores made up of rock with very small or microscopic particles of gold. This gold ore is often found together with quartz or sulfide minerals such as Fool's Gold, which is a pyrite.[26] These are called "lode" deposits. Native gold is also found in the form of free flakes, grains or larger nuggets that have been eroded from rocks and end up in alluvial deposits (called placer deposits). Such free gold is always richer at the surface of gold-bearing veins owing to the oxidation of accompanying minerals followed by weathering, and washing of the dust into streams and rivers, where it collects and can be welded by water action to form nuggets.
Gold sometimes occurs combined with tellurium as the minerals calaverite, krennerite, nagyagite, petzite and sylvanite, and as the rare bismuthide maldonite (Au2Bi) and antimonide aurostibite (AuSb2). Gold also occurs in rare alloys with copper, lead, and mercury: the minerals auricupride (Cu3Au), novodneprite (AuPb3) and weishanite ((Au,Ag)3Hg2). Recent research suggests that microbes can sometimes play an important role in forming gold deposits, transporting and precipitating gold to form grains and nuggets that collect in alluvial deposits.[27]
Isotopes Gold has only one stable isotope, 197Au, which is also its only naturally-occurring isotope. 36 radioisotopes have been synthesized ranging in atomic mass from 169 to 205. The most stable of these is 195Au with a half-life of 186.1 days. 195Au is also the only isotope to decay by electron capture. The least stable is 171Au, which decays by proton emission with a half-life of 30 µs. Most of gold's radioisotopes with atomic masses below 197 decay by some combination of proton emission, α decay, and β+ decay. The exceptions are 195Au, which decays by electron capture, and 196Au, which has a minor β- decay path. All of gold's radioisotopes with atomic masses above 197 decay by β- decay.[28] At least 32 nuclear isomers have also been characterized, ranging in atomic mass from 170 to 200. Within that range, only 178Au, 180Au, 181Au, 182Au, and 188Au do not have isomers. Gold's most stable isomer is 198 m2Au with a half-life of 2.27 days. Gold's least stable isomer is 177 m2Au with a half-life of only 7 ns. 184 m1Au has three decay paths: β+ decay, isomeric transition, and alpha decay. No other isomer or isotope of gold has three decay paths.[28]
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14
Production
The entrance to an underground gold mine in Victoria, Australia
Gold extraction is most economical in large, easily mined deposits. Ore grades as little as 0.5 g/1000 kg (0.5 parts per million, ppm) can be economical. Typical ore grades in open-pit mines are 1–5 g/1000 kg (1–5 ppm); ore grades in underground or hard rock mines are usually at least 3 g/1000 kg (3 ppm). Because ore grades of 30 g/1000 kg (30 ppm) are usually needed before gold is visible to the naked eye, in most gold mines the gold is invisible. Since the 1880s, South Africa has been the
World gold production trend
source for a large proportion of the world’s gold supply, with about 50% of all gold ever produced having come from South Africa. Production in 1970 accounted for 79% of the world supply, producing about 1,000 tonnes. However by 2007 production was just 272 tonnes. This sharp decline was due to the increasing difficulty of extraction, changing economic factors affecting the industry, and tightened safety auditing. In 2007 China (with 276 tonnes) overtook South Africa as the world's largest gold producer, the first time since 1905 that South Africa has not been the largest.[29] The city of Johannesburg located in South
Africa was founded as a result of the Witwatersrand Gold Rush which resulted in the discovery of some of the largest gold deposits the world has ever seen. Gold Gold output in 2005 fields located within the basin in the Free State and Gauteng provinces are extensive in strike and dip requiring some of the world's deepest mines, with the Savuka and TauTona mines being currently the world's deepest gold mine at 3,777 m. The Second Boer War of 1899–1901 between the British Empire and the Afrikaner Boers was at least partly over the rights of miners and possession of the gold wealth in South Africa. Other major producers are the United States, Australia, Russia and Peru. Mines in South Dakota and Nevada supply two-thirds of gold used in the United States. In South America, the controversial project Pascua Lama aims at exploitation of rich fields in the high mountains of Atacama Desert, at the border between Chile and Argentina. Today about one-quarter of the world gold output is estimated to originate from artisanal or small scale mining.[30]
Gold After initial production, gold is often subsequently refined industrially by the Wohlwill process or the Miller process. Other methods of assaying and purifying smaller amounts of gold include parting and inquartation as well as cuppelation, or refining methods based on the dissolution of gold in aqua regia. The world's oceans hold a vast amount of gold, but in very low concentrations (perhaps 1–2 parts per 10 billion, e.g. every cubic kilometer of water could contain 10 to 20 kg of gold). A number of people have claimed to be able to economically recover gold from sea water, but so far they have all been either mistaken or crooks. Reverend Prescott Jernegan ran a gold-from-seawater swindle in the United States in the 1890s. A British fraudster ran the same scam in England in the early 1900s.[31] Fritz Haber (the German inventor of the Haber process) attempted commercial extraction of gold from sea water in an effort to help pay Germany's reparations following World War I. Unfortunately, his assessment of the concentration of gold in sea water was unduly high, probably due to sample contamination. The effort produced little gold and cost the German government far more than the commercial value of the gold recovered. No commercially viable mechanism for performing gold extraction from sea water has yet been identified. Gold synthesis is not economically viable and is unlikely to become so in the foreseeable future. The average gold mining and extraction costs are $238 per troy ounce but these can vary widely depending on mining type and ore quality. In 2001, global mine production amounted to 2,604 tonnes, or 67% of total gold demand in that year. At the end of 2006, it was estimated that all the gold ever mined totaled 158,000 tonnes.[1] This can be represented by a cube with an edge length of just 20.2 meters. Gold is so stable and so valuable that it is always recovered and recycled. There is no true "consumption" of gold in the economic sense; the stock of gold remains essentially constant while ownership shifts from one party to another.[32]
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Consumption India is the world’s largest consumer of gold. Indian consumers buy about 25 per cent of the world’s gold.[33] Indians buy approximately 800 tonnes of gold every year. India is also the largest importer of the yellow metal; in 2008 India imported around 400 tonnes of gold.[34]
Price Like other precious metals, gold is measured by troy weight and by grams. When it is alloyed with other metals the term carat or karat is used to indicate the amount of gold present, with 24 carats being pure gold and lower ratings proportionally less. The purity of a gold bar can also be expressed as a decimal figure ranging from 0 to 1, known as the millesimal fineness, such as 0.995 being very pure. The price of gold is determined on the open market, but a procedure known as the Gold Fixing in London, originating in September 1919, provides a daily benchmark figure to the industry. The afternoon fixing appeared in 1968 to fix a price when US markets are open.
LBMA USD morning price fixings ($US per troy ounce) from Jan 2001 to Apr 2006
Historically gold coinage was widely used as currency; When paper money was introduced, it typically was a receipt redeemable for gold coin or bullion. In an economic system known as the gold standard, a certain weight of gold was given the name of a unit of currency. For a Gold price per ounce in USD since 1968, in actual long period, the United States government US$ and 2006 US$. set the value of the US dollar so that one troy ounce was equal to $20.67 ($664.56/kg), but in 1934 the dollar was devalued to $35.00 per troy ounce ($1125.27/kg). By 1961 it was becoming hard to maintain this price, and a pool of US and European banks agreed to manipulate the market to prevent further currency devaluation against increased gold demand. On March 17, 1968, economic circumstances caused the collapse of the gold pool, and a two-tiered pricing scheme was established whereby gold was still used to settle international accounts at the old $35.00 per troy ounce ($1.13/g) but the price of gold on the private market was allowed to fluctuate; this two-tiered pricing system was abandoned in 1975 when the price of gold was left to find its free-market level. Central banks still hold historical gold reserves as a store of value although the level has generally been declining.
Gold The largest gold depository in the world is that of the U.S. Federal Reserve Bank in New York, which holds about 3% of the gold ever mined, as does the similarly-laden U.S. Bullion Depository at Fort Knox. In 2005 the World Gold Council estimated total global gold supply to be 3,859 tonnes and demand to be 3,754 tonnes, giving a surplus of 105 tonnes.[35]
Price records Since 1968 the price of gold on the open market has ranged widely, from a high of $850/oz ($27,300/kg) on January 21, 1980, to a low of $252.90/oz ($8,131/kg) on June 21, 1999 (London Gold Fixing).[36] The 1980 high was not overtaken until January 3, 2008 when a new maximum of $865.35 per troy ounce was set (a.m. London Gold Fixing).[37] The current record price was set on March 17, 2008 at $1023.50/oz ($32,900/kg)(am. London Gold Fixing).[37]
Long term price trends Since April 2001 the gold price has more than tripled in value against the US dollar,[38] prompting speculation that this long secular bear market (or the Great Commodities Depression) has ended and a bull market has returned.[39] In March 2008, the gold price increased above $1000,[40] which in real terms is still well below the $850/oz. peak on January 21, 1980. Indexed for inflation, the 1980 high would equate to a price of around $2400 in 2007 US dollars. In the last century, major economic crises (such as the Great Depression, World War II, the first and second oil crisis) lowered the Dow/Gold ratio (which is inherently inflation adjusted) substantially, in most cases to a value well below 4.[41] During these difficult times, investors tried to preserve their assets by investing in precious metals, most notably gold and silver.
Compounds Although gold is a noble metal, it forms many and diverse compounds. The oxidation state of gold in its compound ranges from −1 to +5 but Au(I) and Au(III) dominate. Gold(I), referred to as the aurous ion, is the most common oxidation state with “soft” ligands such as thioethers, thiolates, and tertiary phosphines. Au(I) compounds are typically linear. A good example is Au(CN)2−, which is the soluble form of gold encountered in mining. Curiously, aurous complexes of water are rare. The binary gold halides, such as AuCl, form zig-zag polymeric chains, again featuring linear coordination at Au. Most drugs based on gold are Au(I) derivatives.[42] Gold(III) (“auric”) is a common oxidation state and is illustrated by gold(III) chloride, AuCl3. Its derivative is chloroauric acid, HAuCl4, which forms when Au dissolves in aqua regia. Au(III) complexes, like other d8 compounds, are typically square planar.
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Less common oxidation states: Au(-I), Au(II), and Au(V) Compounds containing the Au− anion are called aurides. Caesium auride, CsAu which crystallizes in the caesium chloride motif.[43] Other aurides include those of Rb+, K+, and tetramethylammonium (CH3)4N+.[44] Gold(II) compounds are usually diamagnetic with Au-Au bonds such as [Au(CH2)2P(C6H5)2]2Cl2. A noteworthy, legitimate Au(II) complex is the tetraxenonogold(II) cation, which contains xenon as a ligand, [AuXe4](Sb2F11)2.[45] Gold pentafluoride is the sole example of Au(V), the highest verified oxidation state.[46] Some gold compounds exhibit aurophilic bonding, which describes the tendency of gold ions to interact at distances that are too long to be a conventional Au-Au bond but shorter that van der Waals bonding. The interaction is estimated to be comparable in strength to that of a hydrogen bond.
Mixed valence compounds Well-defined cluster compounds are numerous.[44] In such cases, gold has a fractional oxidation state. A representative example is the octahedral species {Au(P(C6H5)3)}62+. Gold chalcogenides, e.g. "AuS" feature equal amounts of Au(I) and Au(III).
Symbolism Gold has been associated with the extremities of utmost
Swiss-cast 1 kg gold bar
evil and great sanctity throughout history. In the Book of Exodus, the Golden Calf is a symbol of idolatry and rebellion against God. In popular culture, the golden pocket watch and its fastening golden chain were the characteristic accessories of the capitalists, the rich and the industrial tycoons. Credit card companies associate their product with wealth by naming and coloring their top-of-the-range cards “gold” although, in an attempt to out-do each other, platinum has now overtaken gold. In the Book of Genesis, Abraham was said to be rich in gold and silver, and Moses was instructed to cover the Mercy Seat of the Ark of the Covenant with pure gold. Eminent orators such as John Chrysostom were said to have a “mouth of gold with a silver tongue.” Gold is associated with notable anniversaries, particularly in a 50-year cycle, such as a golden wedding anniversary, golden jubilee, etc.
Great human achievements are frequently rewarded with gold, in the form of medals and decorations. Winners of races and prizes are usually awarded the gold medal (such as the Olympic Games and the Nobel Prize), while many award statues are depicted in gold (such as the Academy Awards, the Golden Globe Awards the Emmy Awards, the Palme d'Or, and the British Academy Film Awards). Gold bars at the Emperor Casino in Macau
Medieval kings were inaugurated under the signs of sacred oil and a golden crown, the latter symbolizing the eternal shining light of heaven and thus a Christian king's divinely
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inspired authority. Wedding rings are traditionally made of gold; since it is long-lasting and unaffected by the passage of time, it is considered a suitable material for everyday wear as well as a metaphor for the relationship. In Orthodox Christianity, the wedded couple is adorned with a golden crown during the ceremony, an amalgamation of symbolic rites.
Toxicity Pure gold is non-toxic and non-irritating when ingested[47] and is sometimes used as a food decoration in the form of gold leaf. It is also a component of the alcoholic drinks Goldschläger, Gold Strike, and Goldwasser. Gold is approved as a food additive in the EU (E175 in the Codex Alimentarius). Soluble compounds (gold salts) such as potassium gold cyanide, used in gold electroplating, are toxic to the liver and kidneys. There are rare cases of lethal gold poisoning from potassium gold cyanide.[48] [49] Gold toxicity can be ameliorated with chelation therapy with an agent such as Dimercaprol. It was voted Allergen of the Year in 2001 by the American Contact Dermatitis Society.
See also • • • • • • • • •
Altai Mountains ChipGold Commodity fetishism Digital gold currency Gold fingerprinting Gold Prospectors Association of America Mining in Roman Britain Prospecting Roman engineering
Bibliography • Faulk W, Taylor G (1979) An Immunocolloid Method for the Electron Microscope Immunochemistry 8, 1081–1083. • Kodak (2006) Toning black-and-white materials [50]. Technical Data/Reference sheet G-23, May 2006. • Roth J, Bendayan M, Orci L (1980) FITC-Protein A-Gold Complex for Light and Electron Microscopic Immunocytochemistry. Journal of Histochemistry and Cytochemistry 28, 55–57. • World Gold Council, Jewellery Technology, Jewellery Alloys [51]
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External links • • • •
Getting Gold 1898 book [52] Technical Document on Extraction and Mining of Gold [53] Picture in the Element collection from Heinrich Pniok [54] WebElements.com — Gold [55]
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[27] " Environment & Nature News - Bugs grow gold that looks like coral - 28/01/2004 (http:/ / www. abc. net. au/ science/ news/ enviro/ EnviroRepublish_1032376. htm)". . Retrieved 2006-07-22. This is doctoral research undertaken by Frank Reith at the Australian National University, published 2004. [28] Audi, G. (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3–128. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [29] Mandaro, Laura (2008-01-17). " China now world's largest gold producer; foreign miners at door MarketWatch (http:/ / www. marketwatch. com/ news/ story/ china-now-worlds-largest-gold/ story. aspx?guid={8C528CE8-0262-485D-ACEB-2247D18282CB})". MarketWatch value > research & statistics > statistics > supply and demand statistics (http:/ / www. gold. org/ value/ stats/ statistics/ gold_demand/ index. html)". . Retrieved 2006-07-22. [36] " kitco.com: GOLD - London PM Fix 1975 - present (GIF) (http:/ / kitco. com/ LFgif/ au75-pres. gif)". . Retrieved 2006-07-22. [37] " LBMA statistics (http:/ / www. lbma. org. uk/ 2008dailygold. htm)". Lbma.org.uk. 2008-12-31. . Retrieved 2009-04-05. [38] " kitco.com/LFgif/au3650nyb.gif (http:/ / kitco. com/ LFgif/ au3650nyb. gif)". . [39] " Gold starts 2006 well, but this is not a 25-year high! | Financial Planning (http:/ / www. ameinfo. com/ 75511. html)". Ameinfo.com. . Retrieved 2009-04-05. [40] " 2008 London Gold Fixings (http:/ / www. lbma. org. uk/ 2008dailygold. htm)". Lbma.org.uk. 2008-12-31. . Retrieved 2009-04-05. [41] " upload.wikimedia.org/wikipedia/en/8/84/Longtermdowgoldlogtr1800.png (http:/ / upload. wikimedia. org/ wikipedia/ en/ 8/ 84/ Longtermdowgoldlogtr1800. png)". . [42] Shaw III, C. F. (1999). "Gold-Based Medicinal Agents". Chemical Reviews 99 (9): 2589–2600. doi: 10.1021/cr980431o (http:/ / dx. doi. org/ 10. 1021/ cr980431o). [43] Jansen, Martin (2005). "Effects of relativistic motion of electrons on the chemistry of gold and platinum". Solid State Sciences 7 (12): 1464–1474. doi: 10.1016/j.solidstatesciences.2005.06.015 (http:/ / dx. doi. org/ 10. 1016/ j. solidstatesciences. 2005. 06. 015). [44] Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5. [45] Seidel, S.; Seppelt, K. (2000). "Xenon as a Complex Ligand: The Tetra Xenono Gold(II) Cation in AuXe42+(Sb2F11−)2". Science 290 (5489): 117–118. doi: 10.1126/science.290.5489.117 (http:/ / dx. doi. org/ 10. 1126/ science. 290. 5489. 117). PMID 11021792. [46] Riedel, S.; Kaupp, M. (2006). "Revising the Highest Oxidation States of the 5d Elements: The Case of Iridium(+VII)". Angewandte Chemie International Edition 45 (22): 3708–3711. doi: 10.1002/anie.200600274 (http:/ / dx. doi. org/ 10. 1002/ anie. 200600274). [47] S Dierks (May 2005). " Gold MSDS (http:/ / www. espi-metals. com/ msds's/ gold. htm)". Electronic Space Products International. . [48] I. H. Wright, C. J. Vesey (1986). "Acute poisoning with gold cyanide". Anaesthesia 41 (79): 936–939. doi: 10.1111/j.1365-2044.1986.tb12920.x (http:/ / dx. doi. org/ 10. 1111/ j. 1365-2044. 1986. tb12920. x). [49] Wu, Ming-Ling; Tsai, Wei-Jen; Ger, Jiin; Deng, Jou-Fang; Tsay, Shyh-Haw; Yang, Mo-Hsiung. (2001). "Cholestatic Hepatitis Caused by Acute Gold Potassium Cyanide Poisoning". Clinical toxicology 39 (7): 739–743. doi: 10.1081/CLT-100108516 (http:/ / dx. doi. org/ 10. 1081/ CLT-100108516). [50] [51] [52] [53]
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Image Sources, Licenses and Contributors
Image Sources, Licenses and Contributors image:Au-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Au-TableImage.png License: GNU Free Documentation License Contributors: Conscious, Emilfaro, Korg, Paddy, Saperaud, WeFt, 2 anonymous edits Image: Gold-crystals.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Gold-crystals.jpg License: unknown Contributors: User:Alchemist-hp File:Native_gold_nuggets.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Native_gold_nuggets.jpg License: unknown Contributors: User:Aramgutang File:Nugsrandt.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Nugsrandt.jpg License: Public Domain Contributors: Original uploader was Gold Guru at en.wikipedia File:Ag-Au-Cu-colours-english.svg Source: http://en.wikipedia.org/w/index.php?title=File:Ag-Au-Cu-colours-english.svg License: unknown Contributors: User:Metallos File:Maple Leaf 99999 Gold 2008 Limited ReverseSide.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Maple_Leaf_99999_Gold_2008_Limited_ReverseSide.jpg License: GNU Free Documentation License Contributors: Sloth monkey, 1 anonymous edits File:MocheGoldNecklace.jpg Source: http://en.wikipedia.org/w/index.php?title=File:MocheGoldNecklace.jpg License: unknown Contributors: User:Pattych File:220kg Gold brick Taiwan Museum.jpg Source: http://en.wikipedia.org/w/index.php?title=File:220kg_Gold_brick_Taiwan_Museum.jpg License: Public Domain Contributors: Original uploader was Texcoco at en.wikipedia File:Toi 250kg gold bar.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Toi_250kg_gold_bar.jpg License: unknown Contributors: User:PHGCOM File:Small gold nugget 5mm dia and corresponding foil surface of half sq meter.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Small_gold_nugget_5mm_dia_and_corresponding_foil_surface_of_half_sq_meter.jpg License: unknown Contributors: User:PHGCOM File:Sun symbol.svg Source: http://en.wikipedia.org/w/index.php?title=File:Sun_symbol.svg License: Public Domain Contributors: AnonMoos, Aquantrum, Bvs-aca, Er Komandante, Fibonacci, Juiced lemon, Liftarn, Luccas, Melian, Nagy, Roomba, Rursus, Samuel Grant, SkyBon, Wst, 13 anonymous edits File:TurinPapyrus1.jpg Source: http://en.wikipedia.org/w/index.php?title=File:TurinPapyrus1.jpg License: Public Domain Contributors: Zyzzy File:Tuthankhamun Egyptian Museum.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Tuthankhamun_Egyptian_Museum.jpg License: unknown Contributors: User:Uspn File:Jason Pelias Louvre K127.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Jason_Pelias_Louvre_K127.jpg License: Public Domain Contributors: User:Jastrow File:GoldOreUSGOV.jpg Source: http://en.wikipedia.org/w/index.php?title=File:GoldOreUSGOV.jpg License: Public Domain Contributors: Breeze, Jurema Oliveira, Kluka, Photohound, Saperaud File:Stringer156 nugget.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Stringer156_nugget.jpg License: Public Domain Contributors: Original uploader was Reno Chris at en.wikipedia File:Gold 30g for a 860kg rock.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Gold_30g_for_a_860kg_rock.jpg License: unknown Contributors: User:PHGCOM File:Gold mine.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Gold_mine.jpg License: unknown Contributors: Fir0002, Saperaud File:Gold - world production trend.svg Source: http://en.wikipedia.org/w/index.php?title=File:Gold_-_world_production_trend.svg License: Public Domain Contributors: User:Leyo File:Gold (mined)2.PNG Source: http://en.wikipedia.org/w/index.php?title=File:Gold_(mined)2.PNG License: Public Domain Contributors: User:Anwar_saadat/bubble_maps_(FAQ) File:LBMA Au AM USD 2001-06042006.png Source: http://en.wikipedia.org/w/index.php?title=File:LBMA_Au_AM_USD_2001-06042006.png License: Public Domain Contributors: Emilfaro, NTK File:Gold price.png Source: http://en.wikipedia.org/w/index.php?title=File:Gold_price.png License: unknown Contributors: Donar Reiskoffer File:Goldkey logo removed.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Goldkey_logo_removed.jpg License: Public Domain Contributors: Swiss Banker File:GoldBarAtGrandEmperorCasino.JPG Source: http://en.wikipedia.org/w/index.php?title=File:GoldBarAtGrandEmperorCasino.JPG License: Public Domain Contributors: User:Photnart
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23
Mercury (element)
1
Mercury (element) gold ← mercury → thallium Cd ↑ Hg ↓ Uub Periodic Table Extended Periodic Table General Name, symbol, number
mercury, Hg, 80
Element category
transition metals
Group, Period, Block
12, 6, d
Appearance
silvery
Standard atomic weight
−1
200.59(2) g·mol
Electron configuration
14
[Xe] 4f
Electrons per shell
10
5d
2
6s
2, 8, 18, 32, 18, 2 Physical properties
Phase
liquid Density (near r.t.)
(liquid) 13.534 g·cm−3
Melting point
234.32 K (-38.83 °C, -37.89 °F) Boiling point
629.88 K (356.73 °C, 674.11 °F)
Critical point
1750 K, 172.00 MPa
Heat of fusion
2.29 kJ·mol−1
Heat of vaporization
59.11 kJ·mol−1
Specific heat capacity
(25 °C) 27.983 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
315
350
393
449
523
629
Mercury (element)
2 Atomic properties Crystal structure
rhombohedral
Oxidation states
4, 2 (mercuric), 1 (mercurous) (mildly basic oxide)
Electronegativity
2.00 (Pauling scale)
Ionization energies
1st: 1007.1 kJ/mol 2nd: 1810 kJ/mol 3rd: 3300 kJ/mol
Atomic radius
151 pm
Covalent radius
132±5 pm
Van der Waals radius
155 pm Miscellaneous
Magnetic ordering
diamagnetic
Electrical resistivity
(25 °C) 961 nΩ·m
Thermal conductivity
(300 K) 8.30 W·m
Thermal expansion
(25 °C) 60.4 µm·m
Speed of sound
(liquid, 20 °C) 1451.4 m/s
−1
−1
CAS registry number
−1
·K
−1
·K
7439-97-6 Most-stable isotopes
Main article: Isotopes of mercury iso
NA
half-life
DM
DE (MeV)
DP
194
syn
444 y
ε
0.040
194
195
syn
9.9 h
ε
1.510
195
0.600
197
0.492
203
Hg Hg
196
Hg
197
Hg
0.15% syn
Hg is stable with 116 neutron
64.14 h
9.97%
198
199
16.87%
199
200
23.1%
200
201
13.18%
201
202
29.86%
202
Hg Hg Hg Hg
203
Hg
204
Hg
syn 6.87%
Au
196
198
Hg
Au
ε
Au
Hg is stable with 118 neutron Hg is stable with 119 neutron Hg is stable with 120 neutron Hg is stable with 121 neutron Hg is stable with 122 neutron
46.612 d
β−
Tl
204
Hg is stable with 124 neutron References
Mercury (pronounced /ˈmɜrkjʊri/), also called quicksilver (/ˈkwɪksɪlvər/) or hydrargyrum (/haɪˈdrɑrdʒɨrəm/), is a chemical element with the symbol Hg (Latinized Greek:
Mercury (element)
3
hydrargyrum, meaning watery or liquid silver) and atomic number 80. A heavy, silvery d-block metal, mercury is one of six chemical elements that are liquid at or near room temperature and pressure,[1] [2] the others being caesium, francium, gallium, bromine, and rubidium. Mercury is the only metal that is liquid at standard conditions for temperature and pressure. With a melting point of −38.83 °C and boiling point of 356.73 °C, mercury has one of the widest ranges of its liquid state of any metal. Mercury occurs in deposits throughout the world mostly as cinnabar (mercuric sulfide), which source of the red pigment vermilion, and is mostly obtained by reduction from cinnabar. Cinnabar is highly toxic by ingestion or inhalation of the dust, and mercury poisoning can also result from exposure to soluble forms (such as mercuric chloride or methylmercury), inhalation of mercury vapor, or eating fish contaminated with mercury. Mercury is used in thermometers, barometers, manometers, sphygmomanometers, float valves, and other scientific apparatus, though concerns about the element's toxicity have led to mercury thermometers and sphygmomanometers being largely phased out in clinical environments in favor of alcohol-filled, digital, or thermistor-based instruments. It remains in use in a number of other ways in scientific and scientific research applications, and in amalgam material for dental restoration. It is used in lighting; electricity passed through mercury vapor in a phosphor tube produces short-wave ultraviolet light which then causes the phosphor to fluoresce, making visible light.
Properties Physical Mercury is a heavy, silvery-white metal. As compared to other metals, it is a poor conductor of heat, but a fair conductor of electricity. [3] Mercury has an exceptionally low melting temperature for a d-block metal. A complete explanation of this fact requires a deep excurse into quantum physics, but it can be summarized as follows: mercury has a unique electronic configuration where electrons fill up all the available 1s, 2s, 2p, 3s, 3p, 3d, 4s, 4p, 4d, 4f, 5s, 5p, 5d and 6s Liquid mercury shells. As such configuration strongly resists removal of an electron, mercury behaves similarly to noble gas elements, which form weakly bonded and thus easily melting solids. The stability of the 6s shell is due to the presence of a filled 4f shell. An f shell poorly screens the nuclear charge that increases the attractive Coulomb interaction of the 6s shell and the nucleus (see lanthanide contraction). The absence of a filled inner f shell is the reason for much higher melting temperature of cadmium. Metals such as gold has atoms with one less 6s electron than mercury. Those electrons are more easily removed and are shared between the gold atoms forming relatively strong metallic bonds. [2] [4]
Mercury (element)
4
Reactivity and compounds Mercury dissolves to form amalgams with gold, zinc and many other metals. Because iron is an exception, iron flasks have been traditionally used to trade mercury. Other metals that do not form amalgams with mercury include tantalum, tungsten and platinum. When heated, mercury also reacts with oxygen in air to form mercury oxide, which then can be decomposed by further heating to higher temperatures.[5] Since it is below hydrogen in the reactivity series of metals, mercury does not react with most acids, such as dilute sulfuric acid, though oxidizing acids such as concentrated sulfuric acid and nitric acid or aqua regia dissolve it to give sulfate, nitrate, and chloride salts. Like silver, mercury reacts with atmospheric hydrogen sulfide. Mercury even reacts with solid sulfur flakes, which are used in mercury spill kits to absorb mercury vapors (spill kits also use activated charcoal and powdered zinc).[5] Some important mercury salts include: • Mercury(I) chloride (calomel) is sometimes still used in medicine, acousto-optical filters and as a standard in electrochemistry;[6] • Mercury(II) chloride is a very corrosive, easily sublimating and poisonous substance;[3] • Mercury fulminate, (a detonator widely used in explosives);[3] • Mercury(II) oxide, the main oxide of mercury; • Mercury(II) sulfide (found naturally as the ore cinnabar, or vermilion which is a high-grade paint pigment);[3] • Mercury(II) selenide, Mercury(II) telluride, Mercury cadmium telluride and mercury zinc telluride are semiconductors and infrared detector materials.[7] In these compounds, mercury displays two oxidation states: +1 and +2. The +1 state oxidation involves the dimeric cation, Hg2+2. Solutions of Hg2+2 are in equilibrium with Hg2+ and metallic mercury: Hg2+ + Hg
Hg2+2
This equilibrium causes solutions of Hg2+2 to have a small amount of Hg2+ present. Consuming the Hg2+ by another reaction, such as complexation with strong ligands or precipitation of an insoluble salt, will cause all the Hg2+2 to fully disproportionate to Hg2+ and elemental mercury.[8] Higher oxidation states of mercury were confirmed in September 2007, with the synthesis of mercury(IV) fluoride (HgF4) using matrix isolation techniques.[9] Laboratory tests have found that an electrical discharge causes the noble gases to combine with mercury vapor. These compounds are held together with van der Waals forces and result in Hg·Ne, Hg·Ar, Hg·Kr, and Hg·Xe (see exciplex). Organic mercury compounds are also important. Methylmercury is a dangerous compound that is widely found as a pollutant in water bodies and streams.[10]
Mercury and aluminium Mercury readily combines with aluminium to form a mercury-aluminium amalgam when the two pure metals come into contact. However, when the amalgam is exposed to air, the aluminium oxidizes, leaving mercury behind. The oxide flakes away, exposing more mercury amalgam, which repeats the process. This process continues until the supply of amalgam is exhausted. Because this process releases mercury, a small amount of mercury can "eat through" a large amount of aluminium over time, by progressively forming amalgam and
Mercury (element)
5
relinquishing the aluminium as oxide.[11] Aluminium in air is ordinarily protected by a molecule-thin layer of its own oxide, which is not porous to oxygen. Mercury coming into contact with this oxide does no harm. However, if any elemental aluminium is exposed (even by a recent scratch), the mercury may combine with it, starting the process described above, and potentially damaging a large part of the aluminium before it finally ends.[11] [12] For this reason, restrictions are placed on the use and handling of mercury in proximity with aluminium. In particular, mercury is not allowed aboard aircraft under most circumstances because of the risk of it forming an amalgam with exposed aluminium parts in the aircraft.[11]
Isotopes There are seven stable isotopes of mercury with 202Hg being the most abundant (29.86%). The longest-lived radioisotopes are 194Hg with a half-life of 444 years, and 203Hg with a half-life of 46.612 days. Most of the remaining radioisotopes have half-lives that are less than a day. 199Hg and 201Hg are the most often studied NMR-active nuclei, having spins of 1 ⁄2 and 3⁄2 respectively.[3]
History Mercury was known to the ancient Chinese[13] and was found in Egyptian tombs that date from 1500 BCE.[14] In China and Tibet, mercury use was thought to prolong life, heal fractures, and maintain generally good health. One of China's emperors, Qín Shǐ Huáng Dì — allegedly buried in a tomb that contained rivers of flowing mercury on a model of the land he ruled, representative of the rivers of China — was killed by drinking a mercury and powdered jade mixture (causing liver failure, poisoning, and brain death) intended to give him eternal life.[15] [16] The ancient Greeks used mercury in ointments; the ancient Egyptians and the The symbol for the planet Mercury (☿) Romans used it in cosmetics which sometimes has been used since ancient times to represent the element deformed the face. By 500 BC mercury was used to make amalgams with other metals.[17] The Indian word for alchemy is Rasavātam which means "the way of mercury".[18] Alchemists thought of mercury as the First Matter from which all metals were formed. They believed that different metals could be produced by varying the quality and quantity of sulfur contained within the mercury. The purest of these was gold, and mercury was called for in attempts at the transmutation of base (or impure) metals into gold was the goal of many alchemists.[19] Hg is the modern chemical symbol for mercury. It comes from hydrargyrum, a Latinized form of the Greek word Ύδραργυρος (hydrargyros), which is a compound word meaning "water" and "silver" — since it is liquid, like water, and yet has a silvery metallic sheen. The element was named after the Roman god Mercury, known for speed and mobility. It is associated with the planet Mercury; the astrological symbol for the planet is also one of the alchemical symbols for the metal. Mercury is the only metal for which the alchemical
Mercury (element) planetary name became the common name.[19] The mines in Almadén (Spain), Monte Amiata (Italy), and Idrija (now Slovenia) dominated the mercury production from the opening of the mine in Almadén 2500 years ago until new deposits were found at the end of the 19th century.[20]
Occurrence Mercury is an extremely rare element in the Earth's crust, having an average crustal abundance by mass of only 0.08 parts per million (ppm).[21] However, because it does not blend geochemically with those elements that constitute the majority of the crustal mass, Mercury output in 2005 mercury ores can be extraordinarily concentrated considering the element's abundance in ordinary rock. The richest mercury ores contain up to 2.5% mercury by mass, and even the leanest concentrated deposits are at least 0.1% mercury (12,000 times average crustal abundance). It is found either as a native metal (rare) or in cinnabar, corderoite, livingstonite and other minerals, with cinnabar (HgS) being the most common ore. Mercury ores usually occur in very young orogenic belts where rock of high density are forced to the crust of the Earth, often in hot springs or other volcanic regions.[22] Beginning in 1558, with the invention of the patio process to extract silver from ore using mercury, mercury became an essential resource in the economy of Spain and its American colonies. Mercury was used to extract silver from the lucrative mines in New Spain and Peru. Initially, the Spanish Crown's mines in Almaden in Southern Spain supplied all the mercury for the colonies.[23] Mercury deposits were discovered in the New World, and more than 100,000 tons of mercury were mined from the region of Huancavelica, Peru, over the course of three centuries following the discovery of deposits there in 1563. The patio process and later pan amalgamation process continued to create great demand for mercury to treat silver ores until the late 1800s.[24] Former mines in Italy, the United States and Mexico which once produced a large proportion of the world supply have now been completely mined out or, in the case of Slovenia (Idrija) and Spain (Almadén), shut down due to the fall of the price of mercury. Nevada's McDermitt Mine, the last mercury mine in the United States, closed in 1992. The price of mercury has been highly volatile over the years and in 2006 was $650 per 76-pound (34.46 kg) flask.[25] Mercury is extracted by heating cinnabar in a current of air and condensing the vapor. The equation for this extraction is HgS + O2 → Hg + SO2 In 2005, China was the top producer of mercury with almost two-thirds global share followed by Kyrgyzstan.[26] Several other countries are believed to have unrecorded production of mercury from copper electrowinning processes and by recovery from
6
Mercury (element)
7
effluents. Because of the high toxicity of mercury, both the mining of cinnabar and refining for mercury are hazardous and historic causes of mercury poisoning. In China, prison labor was used by a private mining company as recently as the 1950s to create new cinnabar mercury mines. Thousands of prisoners were used by the Luo Xi mining company to establish new tunnels.[] In addition, worker health in functioning mines is at high risk. The European Union directive calling for compact fluorescent bulbs to be made mandatory by 2012 has encouraged China to re-open deadly cinnabar mines to obtain the mercury required for CFL bulb manufacture. As a result, new generations of Chinese, their livestock, and their crops are being poisoned, particularly in the southern cities of Foshan and Guangzhou, and in the Guizhou province in the southwest.[27] Abandoned mercury mine processing sites often contain very hazardous waste piles of roasted cinnabar calcines. Water runoff from such sites is a recognized source of ecological damage. Former mercury mines may be suited for constructive re-use. For example, in 1976 Santa Clara County, California purchased the historic Almaden Quicksilver Mine and created a county park on the site, after conducting extensive safety and environmental analysis of the property.[28]
Releases in the environment Preindustrial deposition rates of mercury from the atmosphere may be in the range of 4 ng /(1 L of ice deposit). Although that can be considered a natural level of exposure, regional or global sources have significant effects. Volcanic eruptions can increase the atmospheric source by 4 — 6 times.[29] Natural
sources
such
as
volcanoes
are
responsible
for
approximately half of atmospheric mercury emissions. The human-generated half can be divided into the following estimated percentages:[30] [31] [32] • 65% from stationary combustion, of which coal-fired power plants are the largest aggregate source (40% of U.S. mercury emissions in 1999). This includes power plants fueled with gas where the mercury has not been removed. Emissions from coal combustion are between one and two orders of magnitude higher than emissions from oil combustion, depending on the country.[30] • 11% from gold production. The three largest point sources for mercury emissions in the U.S. are the three largest gold mines. • 6.8% from non-ferrous metal production, typically smelters. • 6.4% from cement production.
Amount of atmospheric mercury deposited at Wyoming's Upper Fremont Glacier over the last 270 years
• 3.0% from waste disposal, including municipal and hazardous waste, crematoria, and sewage sludge incineration. This is a significant underestimate due to limited information, and is likely to be off by a factor of two to five. • 3.0% from caustic soda production. • 1.4% from pig iron and steel production.
Mercury (element) • 1.1% from mercury production, mainly for batteries. • 2.0% from other sources. The above percentages are estimates of the global human-caused mercury emissions in 2000, excluding biomass burning, an important source in some regions.[30] Current atmospheric mercury contamination in outdoor urban air is (0.01 – 0.02 µg/m3 ) indoor concentrations are significantly elevated over outdoor concentrations, at a range of 0.0065 – 0.523 µg/m3 (average 0.069 µg/m3) [33] Mercury also enters into the environment through the disposal (e.g., land filling, incineration) of certain products. Products containing mercury include: auto parts, batteries, fluorescent bulbs, medical products, thermometers, and thermostats.[34] Due to health concerns (see below), toxics use reduction efforts are cutting back or eliminating mercury in such products. For example, most thermometers now use pigmented alcohol instead of mercury. Mercury thermometers are still occasionally used in the medical field because they are more accurate than alcohol thermometers, though both are being replaced by electronic thermometers. Mercury thermometers are still widely used for certain scientific applications because of their greater accuracy and working range. The United States Clean Air Act, passed in 1990, put mercury on a list of toxic pollutants that need to be controlled to the greatest possible extent. Thus, industries that release high concentrations of mercury into the environment agreed to install maximum achievable control technologies (MACT). In March 2005 EPA rule[35] added power plants to the list of sources that should be controlled and a national cap and trade rule was issued. States were given until November 2006 to impose stricter controls, and several States are doing so. The rule was being subjected to legal challenges from several States in 2005 and decision was made in 2008. The Clean Air Mercury Rule was struck down by a Federal Appeals Court on February 8, 2008. The rule was deemed not sufficient to protect the health of persons living near coal-fired power plants. The court opinion cited the negative impact on human health from coal fired power plants' mercury emissions documented in the EPA Study Report to Congress of 1998.[36] Historically, one of the largest releases was from the Colex plant, a lithium-isotope separation plant at Oak Ridge. The plant operated in the 1950s and 1960s. Records are incomplete and unclear, but government commissions have estimated that some two million pounds of mercury are unaccounted for.[37] One of the worst industrial disasters in history was caused by the dumping of mercury compounds into Minamata Bay, Japan. The Chisso Corporation, a fertilizer and later petrochemical company, was found responsible for polluting the bay from 1932 — 1968. It is estimated that over 3,000 people suffered various deformities, severe mercury poisoning symptoms or death from what became known as Minamata disease.[38]
8
Mercury (element)
9
Applications Mercury is used primarily for the manufacture of industrial chemicals or for electrical and electronic applications. It is used in some thermometers, especially ones which are used to measure high temperatures. A still increasing amount is used as gaseous mercury in fluorescent lamps, while most of the other applications are slowly phased out due to health and safety regulations and is in some applications replaced with less toxic but considerably more expensive Galinstan alloy.
The bulb of a mercury-in-glass thermometer
Present use Medicine Mercury
and
its
compounds
have
been
used
in
medicine, although they are much less common today than they once were, now that the toxic effects of mercury and its compounds are more widely understood. The element mercury is an ingredient in dental amalgams. Thiomersal (called Thimerosal in the United States) is an organic compound used as a preservative in vaccines, though this use is in decline.[39] Another mercury compound Merbromin (Mercurochrome) is a topical antiseptic used for minor cuts and scrapes is still in use in some countries.
Amalgam filling
Mercury(I) chloride (also known as calomel or mercurous chloride) has traditionally been used as a diuretic, topical disinfectant, and laxative. Mercury(II) chloride (also known as mercuric chloride or corrosive sublimate) was once used to treat syphilis (along with other mercury compounds), although it is so toxic that sometimes the symptoms of its toxicity were confused with those of the syphilis it was believed to treat.[40] It was also used as a disinfectant. Blue mass, a pill or syrup in which mercury is the main ingredient, was prescribed throughout the 1800s for numerous conditions including constipation, depression, child-bearing and toothaches.[41] In the early 20th century, mercury was administered to children yearly as a laxative and dewormer, and it was used in teething powders for infants. The mercury-containing organohalide merbromin (sometimes sold as Mercurochrome) is still widely used but has been banned in some countries such as the U.S.[42]
Mercury (element)
10
Since the 1930s some vaccines have contained the preservative thiomersal, which is metabolized or degraded to ethyl mercury. Although it was widely speculated that this mercury-based preservative can cause or trigger autism in children, scientific studies showed no evidence supporting any such link.[43] Nevertheless thiomersal has been removed from or reduced to trace amounts in all U.S. vaccines recommended for children 6 years of age and under, with the exception of inactivated influenza vaccine.[44]
The deep violet glow of a mercury vapor discharge in a germicidal lamp, whose spectrum is rich in invisible ultraviolet radiation.
Mercury in the form of one of its common ores, cinnabar, remains an important component of Chinese, Tibetan, and Ayurvedic medicine. As problems may arise when these medicines are exported to countries that prohibit the use of mercury in medicines, in recent times, less toxic substitutes have been devised. Today, the use of mercury in medicine has greatly declined in all respects, especially in developed countries. Thermometers and sphygmomanometers containing mercury were invented in the early 18th and late 19th centuries, respectively. In the early 21st century, their use is declining and has been banned in some countries, states and medical institutions. In 2002, the U.S. Senate passed legislation to phase out the sale of non-prescription mercury thermometers. In 2003, Washington and Maine became the first states to ban mercury blood pressure devices.[45] Mercury compounds are found in some over-the-counter drugs, including topical antiseptics, stimulant laxatives, diaper-rash ointment, eye drops, and nasal sprays. The FDA has “inadequate data to establish general recognition of the safety and effectiveness,” of the mercury ingredients in these products.[46] Mercury is still used in some diuretics, although substitutes now exist for most therapeutic uses.
Mercury (element) Cosmetics Mercury, as thiomersal, is widely used in the manufacture of mascara. In 2008, Minnesota became the first state in the US to ban intentionally added mercury in cosmetics, giving it a tougher standard than the federal government.[47] Production of chlorine and caustic soda Chlorine is produced from sodium chloride (common salt, NaCl) using electrolysis to separate the metallic sodium from the chlorine gas. Usually the salt is dissolved in water to produce a brine. By-products of any such chloralkali process are hydrogen (H2) and sodium hydroxide (NaOH), which is commonly called caustic soda or lye. By far the largest use of mercury[48] [49] in the late 1900s was in the mercury cell process (also called the Castner-Kellner process) where metallic sodium is formed as an amalgam at a cathode made from mercury; this sodium is then reacted with water to produce sodium hydroxide.[50] Many of the industrial mercury releases of the 1900s came from this process, although modern plants claimed to be safe in this regard.[49] After about 1985, all new chloralkali production facilities that were built in the United States used either membrane cell or diaphragm cell technologies to produce chlorine. Gold and silver mining Historically, mercury was used extensively in hydraulic gold mining in order to help the gold to sink through the flowing water-gravel mixture. Thin mercury particles may form mercury-gold amalgam and therefore increase the gold recovery rates.[3] Large scale use of mercury stopped in the 1960s. However, mercury is still used in small scale, often clandestine, gold prospection. It is estimated that 45,000 metric tons of mercury used in California for placer mining have not been recovered.[51] Mercury was also used in silver mining.[52] Other present uses Gaseous mercury is used in mercury-vapor lamps and some "neon sign" type advertising signs and fluorescent lamps. Those low-pressure lamps emit very spectrally narrow lines, which are traditionally used in optical spectroscopy for calibration of spectral position. Commercial calibration lamps are sold for this purpose; however simply reflecting some of the fluorescent-lamp ceiling light into a spectrometer is a common calibration practice.[53] Gaseous mercury is also found Skin tanner containing a low-pressure in some electron tubes, including ignitrons, thyratrons, mercury vapor lamp and two infrared and mercury arc rectifiers.[54] It is also used in lamps, which act both as light source specialty medical care lamps for skin tanning and and electrical ballast [55] desinfection (see pictures). Gaseous mercury is added to cold cathode argon-filled lamps to increase the ionization and electrical conductivity. An argon filled lamp without mercury will have dull spots and will fail to light correctly. Lighting containing mercury can be
11
Mercury (element)
12
bombarded/oven pumped only once. When added to neon filled tubes the light produced will be inconsistent red/blue spots until the initial burning-in process is completed; eventually it will light a consistent dull off-blue color.[56] Some medical thermometers, especially those for high temperatures, are filled with mercury, however, they are gradually disappearing. In the United States, non-prescription sale of mercury fever thermometers Assorted types of fluorescent lamps. has been banned in 2003. [57] It is also found in liquid-mirror telescopes. The mirror is formed by rotating liquid mercury on a disk, the parabolic form of the liquid thus formed reflecting and focusing incident light. Such telescopes are cheaper than conventional large mirror telescopes by up to a factor of 100, but the mirror cannot be tilted and always points straight up.[58] [59] Liquid mercury is a part of popular secondary reference electrode (called the calomel electrode) in electrochemistry as an alternative to the standard hydrogen electrode. The calomel electrode is used to work out the electrode potential of half cells.[60] Last, but not least, the triple point of mercury, -38.8344 °C, is a fixed point used as a temperature standard for the International Temperature Scale (ITS-90).[3]
Proposed uses Liquid mercury has been proposed as a working fluid for a heat pipe type of cooling device for spacecraft heat rejection systems or radiation panels.[61] A new type of atomic clock, using mercury instead of caesium, has been demonstrated. Accuracy is expected to be within one second in 100 million years.[62] [63]
Historic uses Mercury was used for preserving wood, developing daguerreotypes, silvering mirrors, anti-fouling paints (discontinued in 1990), herbicides (discontinued in 1995), handheld maze games, cleaning, and road leveling devices in cars. Mercury compounds have been used in antiseptics, laxatives, antidepressants, and in antisyphilitics. It was also allegedly used by allied spies to sabotage German planes. A mercury paste was applied to bare aluminium, causing the metal to rapidly corrode. This would cause structural failures.[12]
Old mercury switches
• Mercury switches (including home mercury light switches installed prior to 1970), tilt switches used in old fire detectors, tilt switches in many modern home thermostats,[64] electrodes in some types of electrolysis, batteries (mercury cells), sodium hydroxide and chlorine production, handheld games, catalysts, insecticides and liquid mirror telescopes.[65]
Mercury (element)
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• In Islamic Spain it was used for filling decorative pools. Later the American artist Alexander Calder built a mercury fountain for the Spanish Pavilion at the 1937 World Exhibition in Paris. The fountain is now on display at the Fundació Joan Miró in Barcelona.[66] • Mercury was used inside wobbler lures. Its heavy, liquid form made it useful since the lures made an attractive irregular movement when the mercury moved inside the plug. Such use was stopped due to environmental concerns, but illegal preparation of modern fishing plugs has occurred.
Mercury manometer to measure pressure
• The Fresnel lenses of old lighthouses used to float and rotate in a bath of mercury which acted like a bearing.[67] • Mercury sphygmomanometers (blood pressure meter), barometers, diffusion pumps, coulometers, and many other laboratory instruments. As an opaque liquid with a high density and a nearly linear thermal expansion, it is ideal for this role.[68] • Liquid mercury was used as a coolant for some nuclear reactors; however, sodium is proposed for reactors cooled with liquid metal, because the high density of mercury requires much more energy to circulate as coolant.[69] • Mercury was a propellant for early ion engines in electric propulsion systems. Advantages were mercury's high molecular weight, low ionization energy, low dual-ionization energy, high liquid density and liquid storability at room temperature. Disadvantages were concerns regarding environmental impact associated with ground testing and concerns about eventual cooling and condensation of some of the propellant on the spacecraft in long-duration operations. The first spaceflight to use electric propulsion was a mercury-fueled ion thruster SERT-1 launched by NASA at its Wallops Flight Facility in 1964. SERT stands for Space Electric Rocket Test. The SERT-1 flight was followed up by the SERT-2 flight in 1970. Mercury and caesium were preferred propellants for ion engines until Hughes Research Laboratory performed studies finding xenon gas to be a suitable replacement. Xenon is now the preferred propellant for ion engines as it has a high molecular weight, little or no reactivity due its noble gas nature, and has a high liquid density under mild cryogenic storage.[70] [71] • Experimental mercury vapor turbines were installed to increase the efficiency of fossil-fuel electrical power plants.[72] • Mercury was once used as a gun barrel bore cleaner.[73] [74] Hat making From the mid-18th to the mid-19th centuries, a process called "carroting" was used in the making of felt hats. Animal skins were rinsed in an orange solution (the term "carroting" arose from this color) of the mercury compound mercuric nitrate, Hg(NO3)2·2H2O.[75] This process separated the fur from the pelt and matted it together. This solution and the vapors it produced were highly toxic. The United States Public Health Service banned the use of mercury in the felt industry in December 1941. The psychological symptoms associated
Mercury (element) with mercury poisoning are said by some to have inspired the phrase "mad as a hatter", though etymological study suggests that the phrase is actually much older and unrelated to hatters - see hatter for commentary on the origin of the phrase. Lewis Carroll's "Mad Hatter" in his book Alice's Adventures in Wonderland was a play on words based on the [76] older phrase, but the character himself does not exhibit symptoms of mercury poisoning.
Safety Mercury and most of its compounds are extremely toxic and are generally handled with care; in cases of spills involving mercury (such as from certain thermometers or fluorescent light bulbs) specific cleaning procedures are used to avoid toxic exposure.[77] It can be inhaled and absorbed through the skin and mucous membranes, so containers of mercury are securely sealed to avoid spills and evaporation. Heating of mercury, or compounds of mercury that may decompose when heated, are always carried out with adequate ventilation in order to avoid exposure to mercury vapor. The most toxic forms of mercury are its organic compounds, such as dimethylmercury and methylmercury. However, inorganic compounds, such as cinnabar are also highly toxic by ingestion or inhalation of the dust.[78] Mercury can cause both chronic and acute poisoning.
Occupational exposure Due to the health effects of mercury exposure, industrial and commercial uses are regulated in many countries. The World Health Organization, OSHA, and NIOSH all treat mercury as an occupational hazard, and have established specific occupational exposure limits. Environmental releases and disposal of mercury are regulated in the U.S. primarily by the United States Environmental Protection Agency. Case control studies have shown effects such as tremors, impaired cognitive skills, and sleep disturbance in workers with chronic exposure to mercury vapor even at low concentrations in the range 0.7–42 μg/m3.[79] [80] A study has shown that acute exposure (4 — 8 hours) to calculated elemental mercury levels of 1.1 to 44 mg/m3 resulted in chest pain, dyspnea, cough, hemoptysis, impairment of pulmonary function, and evidence of interstitial pneumonitis.[81] Acute exposure to mercury vapor has been shown to result in profound central nervous system effects, including psychotic reactions characterized by delirium, hallucinations, and suicidal tendency. Occupational exposure has resulted in broad-ranging functional disturbance, including erethism, irritability, excitability, excessive shyness, and insomnia. With continuing exposure, a fine tremor develops and may escalate to violent muscular spasms. Tremor initially involves the hands and later spreads to the eyelids, lips, and tongue. Long-term, low-level exposure has been associated with more subtle symptoms of erethism, including fatigue, irritability, loss of memory, vivid dreams, and depression.[82] [83]
Treatment Research on the treatment of mercury poisoning is limited. Currently available drugs for acute mercurial poisoning include chelators N-acetyl-D,L-penicillamine (NAP), British Anti-Lewisite (BAL), 2,3-dimercapto-1-propanesulfonic acid (DMPS), and dimercaptosuccinic acid (DMSA). In one small study including 11 construction workers exposed to elemental mercury, patients were treated with DMSA and NAP.[84] Chelation therapy with both drugs resulted in the mobilization of a small fraction of the total
14
Mercury (element) estimated body mercury. DMSA was able to increase the excretion of mercury to a greater extent than NAP.[85]
Fish Fish and shellfish have a natural tendency to concentrate mercury in their bodies, often in the form of methylmercury, a highly toxic organic compound of mercury. Species of fish that are high on the food chain, such as shark, swordfish, king mackerel, albacore tuna, and tilefish contain higher concentrations of mercury than others. As mercury and methylmercury are fat soluble, they primarily accumulate in the viscera, although they are also found throughout the muscle tissue. When this fish is consumed by a predator, the mercury level is accumulated. Since fish are less efficient at depurating than accumulating methylmercury, fish-tissue concentrations increase over time. Thus species that are high on the food chain amass body burdens of mercury that can be ten times higher than the species they consume. This process is called biomagnification. Mercury poisoning happened this way in Minamata, Japan, now called Minamata disease. As a result, those consuming high levels of fish should be aware of the symptoms of mercury poisoning. [86]
Regulations In the United States, the Environmental Protection Agency is charged with regulating and managing mercury contamination. Several laws give the EPA this authority, including the Clean Air Act, the Clean Water Act, the Resource Conservation and Recovery Act, and the Safe Drinking Water Act. Additionally, the Mercury-Containing and Rechargeable Battery Management Act, passed in 1996, phases out the use of mercury in batteries, and provides for the efficient and cost-effective disposal of many types of used batteries.[87] North America contributed approximately 11% of the total global anthropogenic mercury emissions in 1995.[88] In the European Union, the directive on the Restriction of the Use of Certain Hazardous Substances in Electrical and Electronic Equipment (see RoHS) bans mercury from certain electrical and electronic products, and limits the amount of mercury in other products to less than 1000 ppm.[89] There are restriction for mercury concentration in packaging (the limit is 100 ppm for sum of mercury, lead, hexavalent chromium and cadmium) and batteries (the limit is 5 ppm).[90] In July 2007, the European Union also banned mercury in non-electrical measuring devices, such as thermometers and barometers. The ban applies to new devices only, and contains exemptions for the health care sector and a two year grace period for manufacturers of barometers. [91] Norway enacted a total ban on the use of mercury in the manufacturing and import/export of mercury products, effective January 1, 2008.[92] In 2002, several lakes in Norway were found to have a poor state of mercury pollution, with an excess of 1 mg/g of mercury in their sediment.[93]
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Mercury (element)
Further reading • Jane M. Hightower Diagnosis: Mercury: Money, Politics, and Poison Island Press (October 1, 2008) ISBN 1597263958 ISBN 978-1597263955
External links • • • • • • • • • •
ATSDR — ToxFAQs: Mercury [94] Centers for Disease Control and Prevention - Mercury Topic [95] EPA fish consumption guidelines [96] Global Mercury Assessment report 2002 [97] by the UNEP. Global Mercury Project [98] Got Mercury? calculator [99] Hg 80 Mercury [100] Japanese Sushi Lovers Shrug at Mercury [101] Material Safety Data Sheet — Mercury [102] Mercury Contamination in fish and Source Control, Oceana [103]
• Mercury (UK PID) [104]. National Poisons Information Service: Medical Toxicology Unit (London Centre) - Kolev, S.T. Bates, N. • Natural Resources Defense Council (NRDC): Mercury Contamination in Fish guide [105] — NRDC • NLM Hazardous Substances Databank — Mercury [106] • University of Calgary: How Mercury Causes Brain Neuron Degeneration [107] • WebElements.com — Mercury [108] • Mercury and Your Health [109] pnb:ہراپ
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Mercury (element) [62] " World's most precise clock developed (http:/ / news. bbc. co. uk/ 1/ hi/ sci/ tech/ 1435488. stm)". BBC. 2001-07-12. . Retrieved 2007-05-01. [63] " NIST New Atomic Clock Could Be 1,000 Times Better Than Today’s Best (http:/ / www. nist. gov/ public_affairs/ update/ upd010723. htm#Time)". NIST update. July 23, 2001. NIST. [64] Shelton, C (2004). Electrical Installations (http:/ / books. google. com/ books?id=cxPEiSXh44cC& pg=PA260). Nelson Thornes. p. 260. ISBN 0748779795. . [65] " The Large Zenith Telescope (http:/ / www. astro. ubc. ca/ LMT/ lzt/ index. html)". . Retrieved 2009-07-07. [66] Lew K (2008). Mercury (http:/ / books. google. co. jp/ books?id=pgUfSrD4gzQC& pg=PA10). The Rosen Publishing Group. p. 10. ISBN 1404217800. . [67] Pearson L F (2003). Lighthouses (http:/ / books. google. com/ books?id=oY8nG-6B6v0C& pg=PA29). Osprey Publishing. p. 29. ISBN 0747805563. . [68] Ramanathan E. AIEEE Chemistry (http:/ / books. google. com/ books?id=G8QyI1Nf0VQC& pg=PA251). Sura Books. p. 251. ISBN 8172542933. . [69] Collier (1987). Introduction to Nuclear Power (http:/ / books. google. com/ books?id=2KYVftKE9NUC& pg=PA64). Taylor & Francis. p. 64. ISBN 1560326824. . [70] " Glenn Contributions to Deep Space 1 (http:/ / www. nasa. gov/ centers/ glenn/ about/ history/ ds1. html)". NASA. . Retrieved 2009-07-07. [71] " Electric space propulsion (http:/ / www. daviddarling. info/ encyclopedia/ E/ electricprop. html)". . Retrieved 2009-07-07. [72] Popular Science (http:/ / books. google. com/ books?id=9ycDAAAAMBAJ& pg=PA40). 118, No. 3. Bonnier Corporation. 1931. p. 40. . [73] Francis, G. W. (1849). Chemical Experiments. D. Francis. pp. 62. [74] Castles, WT; Kimball, VF (2005). Firearms and Their Use. Kessinger Publishing. pp. 104. ISBN 9781417989577. [75] Lee, J.D. (1999). Concise Inorganic Chemistry. Wiley-Blackwell. ISBN 9780632052936. [76] Waldron, HA (1983). "Did the Mad Hatter have mercury poisoning?". Br Med J (Clin Res Ed) 287: 1961. doi: 10.1136/bmj.287.6409.1961 (http:/ / dx. doi. org/ 10. 1136/ bmj. 287. 6409. 1961). PMID 6418283. [77] " Mercury: Spills, Disposal and Site Cleanup (http:/ / www. epa. gov/ mercury/ spills/ index. htm)". Environmental Protection Agency. . Retrieved 2007-08-11. [78] " Safety data for mercuric sulphide (http:/ / msds. chem. ox. ac. uk/ ME/ mercuric_sulphide. html)". Oxford University. . Retrieved 2009-07-07. [79] Ngim, CH; Foo, SC; Boey KW; and Keyaratnam J (1992). "Chronic neurobehavioral effects of elemental mercury in dentists". British Journal of Industrial Medicine 49: 782. [80] Liang YX, Sun RK, Chen ZQ, and Li LH (1993). "Psychological effects of low exposure to mercury vapor: Application of computer-administered neurobehavioral evaluation system". Environmental Research 60: 320. doi: 10.1006/enrs.1993.1040 (http:/ / dx. doi. org/ 10. 1006/ enrs. 1993. 1040). [81] McFarland, RB and Reigel, H (1978). "Chronic Mercury Poisoning from a Single Brief Exposure". J. Occup. Med. 20: 532. doi: 10.1097/00043764-197808000-00003 (http:/ / dx. doi. org/ 10. 1097/ 00043764-197808000-00003). [82] Environmental Health Criteria 1: Mercury (http:/ / www. inchem. org/ documents/ ehc/ ehc/ ehc001. htm). Geneva: World Health Organization. 1976. ISBN 9241540613. . [83] published under the joint sponsorship of the United Nations Environment Programme, the International Labour Organisation, and the World Health Organization ; first draft prep. by L. Friberg. (1991). Inorganic mercury. Environmental Health Criteria 118 (http:/ / www. inchem. org/ documents/ ehc/ ehc/ ehc118. htm). Geneva: World Health Organization. ISBN 9241571187. . [84] Bluhm, RE, et al. (1992). "Elemental Mercury Vapour Toxicity, Treatment, and Prognosis After Acute, Intensive Exposure in Chloralkali Plant Workers. Part I: History, Neuropsychological Findings and Chelator effects". Hum Exp Toxicol 11: 201. doi: 10.1177/096032719201100308 (http:/ / dx. doi. org/ 10. 1177/ 096032719201100308). [85] Bluhm, Re; Bobbitt, Rg; Welch, Lw; Wood, Aj; Bonfiglio, Jf; Sarzen, C; Heath, Aj; Branch, Ra (May 1992). "Elemental mercury vapour toxicity, treatment, and prognosis after acute, intensive exposure in chloralkali plant workers. Part I: History, neuropsychological findings and chelator effects.". Human & experimental toxicology 11 (3): 201–10. doi: 10.1177/096032719201100308 (http:/ / dx. doi. org/ 10. 1177/ 096032719201100308). ISSN 0960-3271 (http:/ / worldcat. org/ issn/ 0960-3271). PMID 1352115. [86] Wallace, H.. " Mercury and Your Health (http:/ / www. hannahmwallace. typepad. com/ hannahs_clips/ 2009/ 05/ mercury-your-health-. html)". . [87] " Mercury: Laws and regulations (http:/ / www. epa. gov/ mercury/ regs. htm)". United States Environmental Protection Agency. April 16, 2008. . Retrieved 2008-05-30.
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Mercury (element) [88] " Reductions in Mercury Emissons (http:/ / www. ijc. org/ php/ publications/ html/ 12br/ english/ report/ chemical/ rme. html)". International Joint Commission on the Great Lakes. . [89] " Directive on the Restriction of the Use of Certain Hazardous Substances in Electrical and Electronic Equipment (http:/ / eur-lex. europa. eu/ LexUriServ/ LexUriServ. do?uri=OJ:L:2003:037:0019:0023:EN:PDF)". 2002/95/EC. . Article 4 Paragraph 1. e.g. "Member States shall ensure that, from July 1, 2006, new electrical and electronic equipment put on the market does not contain lead, mercury, cadmium, hexavalent chromium, polybrominated biphenyls (PBB) or polybrominated diphenyl ethers (PBDE)." [90] " Mercury compounds in European Union: (http:/ / www. eiatrack. org/ s/ 1785)". EIA Track (http:/ / www. eiatrack. org/ ). 2007. . Retrieved 2008-05-30. [91] Jones H. (July 10, 2007). " EU bans mercury in barometers, thermometers (http:/ / www. reuters. com/ article/ environmentNews/ idUSL0988544920070710)". Reuters. . Retrieved 2008-05-30. [92] " Norway to ban mercury (http:/ / www. eubusiness. com/ news-eu/ 1198237627. 85)". EU Business (http:/ / www. eubusiness. com/ ). December 21, 2007. . Retrieved 2008-05-30. [93] Berg, T; Fjeld, E; Steinnes, E (Sep 2006). "Atmospheric mercury in Norway: contributions from different sources.". The Science of the total environment 368 (1): 3–9. doi: 10.1016/j.scitotenv.2005.09.059 (http:/ / dx. doi. org/ 10. 1016/ j. scitotenv. 2005. 09. 059). ISSN 0048-9697 (http:/ / worldcat. org/ issn/ 0048-9697). PMID 16310836. [94] [95] [96] [97] [98]
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20
Article Sources and Contributors
Article Sources and Contributors Mercury (element) Source: http://en.wikipedia.org/w/index.php?oldid=309029303 Contributors: (, (jarbarf), 130.94.122.xxx, 216.102.97.xxx, 2D, 2over0, 3Jane, 64.12.105.xxx, 7&6=thirteen, A2Kafir, ABF, AJR, Achim1999, Adambro, Adashiel, Addshore, AdjustShift, Admiral Norton, Afluent Rider, Agateller, Agiorgio, Ahoerstemeier, Aitias, Ajaxkroon, Ajonlime, Alansohn, Ale jrb, Alex.muller, AlexTingle, Alexmcfire, Alksub, Allstarecho, Almadenbuff, Amaraiel, AndonicO, Andre Engels, Andrew Kelly, Andrewpmk, Anetode, Angela, Anil1956, Animum, Anoop.m, Antandrus, Anwar saadat, Apoc2400, Aquahelper, Arcadian, Arch dude, Archimerged, ArglebargleIV, Arock0930, Arthurbrown, Asmodeus Samael, AssegaiAli, Asterion, AtteLynx, Atulsnischal, Audacitor, Auranor, Aussie Alchemist, Avenue, Axiosaurus, BCAttwood, BabuBhatt, Baccyak4H, BartBenjamin, Barticus88, Bass fishing physicist, Bassbonerocks, Beetstra, Bejnar, Beland, Ben Tre, Bencherlite, Bender235, Benimatt, Benjiboi, Betacommand, Bfigura's puppy, Big Bird, Bionerd, BjKa, Blainster, Blanchardb, Bleeisme, BlueEarth, Bobblewik, Bobby H. Heffley, Bobjoejill, Bobo12345, Bobo192, Bogey97, Bongwarrior, Borbrav, BorgQueen, Borisattva, Bovlk, Brian Huffman, BrokenSphere, Bryan Derksen, Bryanm61, Bubba hotep, BullRangifer, Burnedthru, Bwhack, CYD, Cacycle, Cam123cam123, CambridgeBayWeather, Capricorn42, Captain Infinity, Carnildo, Catalystman, Cayzle, Cdang, Centrx, Ceorl, Ceranthor, Chameleon, Cheesybananas, ChicXulub, Chiu frederick, Chowbok, Chris 73, Chrislk02, Chrisrwood, Closedmouth, Cmagnan, Cmichael, Coconuteire, Coffee, Cometstyles, Common Man, ConorHil, Conversion script, Corey Bryant, Cornbinks, CosineKitty, Cosmium, Courtjester555, Cow1995, Crindalli21, Crusadeonilliteracy, Cst17, Cuvtixo, Cyclonenim, D6, DAMurphy, DJ Creamity, DStoykov, DVD R W, DabMachine, Dabomb87, Dadude3320, Dalstadt, Dancter, Daniel 1992, Darrien, Darth Panda, Davehi1, Daven200520, David.cormier, DeadEyeArrow, Deeptrivia, Deglr6328, Deli nk, Deor, DerHexer, Derek.cashman, Dfe6543, Dforest, Dinofinatic, Discospinster, Dlae, Dmoss, Dnvrfantj, 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Image Sources, Licenses and Contributors image:Hg-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Hg-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Conscious, Paddy, Saperaud Image: Hg_Mercury.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Hg_Mercury.jpg License: Public Domain Contributors: Achim Hering, Ies, Kluka, Tano4595, Vonvon, А, 10 anonymous edits Image:Pouring liquid mercury bionerd.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Pouring_liquid_mercury_bionerd.jpg License: unknown Contributors: Own work Image:Equilibrium.svg Source: http://en.wikipedia.org/w/index.php?title=File:Equilibrium.svg License: Public Domain Contributors: User:L'Aquatique Image:Mercury symbol.svg Source: http://en.wikipedia.org/w/index.php?title=File:Mercury_symbol.svg License: Public Domain Contributors: Aquantrum, Badseed, Basilicofresco, Herbythyme, Lexicon, Mrcht, Rursus, Ruslik0, Stanmar, Starwiz, Urhixidur, 4 anonymous edits Image:Mercury output2.PNG Source: http://en.wikipedia.org/w/index.php?title=File:Mercury_output2.PNG License: Public Domain Contributors: Original uploader was Anwar saadat at en.wikipedia
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Image Sources, Licenses and Contributors Image:Mercury fremont ice core.png Source: http://en.wikipedia.org/w/index.php?title=File:Mercury_fremont_ice_core.png License: unknown Contributors: US government Image:Maximum thermometer close up 2.JPG Source: http://en.wikipedia.org/w/index.php?title=File:Maximum_thermometer_close_up_2.JPG License: Public Domain Contributors: CambridgeBayWeather, Christophe.Finot, LimoWreck, Saperaud Image:Amalgam.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Amalgam.jpg License: Public Domain Contributors: w:de:Benutzer:Ulrich BirkhoffUlrich Birkhoff Image:Germicidal UV discharge tube glow.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Germicidal_UV_discharge_tube_glow.jpg License: GNU Free Documentation License Contributors: Deglr6328 Image:Mercuryvaporlamp.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Mercuryvaporlamp.jpg License: Public Domain Contributors: Skatebiker Image:Leuchtstofflampen-chtaube050409.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Leuchtstofflampen-chtaube050409.jpg License: unknown Contributors: User:Chtaube, User:Deglr6328 Image:Old mercury switches bionerd.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Old_mercury_switches_bionerd.jpg License: unknown Contributors: Sathya Image:Barometer mercury column hg.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Barometer_mercury_column_hg.jpg License: unknown Contributors: User:Hgrobe
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
22
Thallium
1
Thallium mercury ← thallium → lead In ↑ Tl ↓ Uut Periodic Table Extended Periodic Table General Name, symbol, number
thallium, Tl, 81
Element category
poor metals
Group, Period, Block
13, 6, p silvery white
Appearance
Standard atomic weight
−1
204.3833(2) g·mol
Electron configuration
14
[Xe] 4f
Electrons per shell
10
5d
2
1
6s 6p
2, 8, 18, 32, 18, 3 Physical properties
Phase
solid Density (near r.t.)
11.85 g·cm−3
Liquid density at m.p.
11.22 g·cm−3
Melting point
577 K (304 °C, 579 °F) Boiling point
1746 K (1473 °C, 2683 °F)
Heat of fusion
4.14 kJ·mol−1
Heat of vaporization
165 kJ·mol−1
Specific heat capacity
(25 °C) 26.32 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
882
977
1097
1252
1461
1758
Thallium
2 Atomic properties Crystal structure
hexagonal
Oxidation states
3, 1 (mildly basic oxide)
Electronegativity
1.62 (Pauling scale)
Ionization energies
1st: 589.4 kJ/mol 2nd: 1971 kJ/mol 3rd: 2878 kJ/mol
Atomic radius
170 pm
Covalent radius
170±8 pm
Van der Waals radius
196 pm Miscellaneous [1]
Magnetic ordering
diamagnetic
Electrical resistivity
(20 °C) 0.18 µ Ω·m
Thermal conductivity
(300 K) 46.1 W·m
Thermal expansion
(25 °C) 29.9 µm·m
Speed of sound (thin rod)
(20 °C) 818 m/s
Young's modulus
8 GPa
Shear modulus
2.8 GPa Bulk modulus
−1
·K
−1
−1
·K
−1
43 GPa
Poisson ratio
0.45
Mohs hardness
1.2 Brinell hardness
26.4 MPa
CAS registry number
7440-28-0 Most-stable isotopes
Main article: Isotopes of thallium iso 203
Tl
204
Tl
205
Tl
NA 29.524% syn
70.476%
half-life
DM
DE (MeV)
DP
203
Tl is stable with 122 neutron
119 Ms (3.78 y)
β−
0.764
204
ε
0.347
204
Pb Hg
205
Tl is stable with 124 neutron References
Thallium (pronounced /ˈθæliəm/) is a chemical element with the symbol Tl and atomic number 81. This soft gray malleable poor metal resembles tin but discolors when exposed to air. Approximately 60-70% of thallium production is used in the electronics industry, and the rest is used in the pharmaceutical industry and in glass manufacturing.[2] It is also used in infrared detectors. Thallium is highly toxic and is used in rat poisons and insecticides,
Thallium but its use has been cut back or eliminated in many countries. Because of its use for murder, thallium has gained the nicknames "The Poisoner's Poison" and "Inheritance Powder" (alongside arsenic).
Characteristics Thallium is very soft and malleable and can be cut with a knife at room temperature. It has a metallic luster, but when exposed to air, it quickly tarnishes with a bluish-grey tinge that resembles lead. (It is preserved by keeping it under oil). A heavy layer of oxide builds up on thallium if left in air. In the presence of water, thallium hydroxide is formed.
History Thallium (Greek θαλλός, thallos, meaning "a green shoot or twig")[3] was discovered by flame spectroscopy in 1862. The name comes from thallium's bright green spectral emission lines. After the publication of the improved method of flame spectroscopy by Robert Bunsen and Gustav Kirchhoff[4] and the discovery of caesium and rubidium in the years 1859 to 1860 flame spectroscopy became an approved method to determine the composition of minerals and chemical products. William Crookes and Claude-Auguste Lamy both started to use the new method. William Crookes used it to make spectroscopic determinations for tellurium on selenium compounds deposited in the lead chamber of a sulfuric acid production plant near Tilkerode in the Harz mountains. He had obtained the samples for his research on selenium cyanide from August Hofmann years earlier.[5] [6] By 1862 Crookes was able to isolate small quantities of the element and determine the properties of a few compounds.[7] Claude-Auguste Lamy used a similar spectrometer to Crookes' to determine the composition of a selenium-containing substance which was deposited during the production of sulfuric acid from pyrite. He also noticed the new green line in the spectra and concluded that a new element was present. Lamy had received this material from the sulfuric acid plant of his friend Fréd Kuhlmann and this by-product was available in large quantities. Lamy started to isolate the new element from that source.[8] The fact that Lamy was able to work ample quantities of thallium enabled him to determine the properties of several compounds and in addition he prepared a small ingot of metallic thallium which he prepared by remelting thallium he had obtained by electrolysis of thallium salts. As both scientists discovered thallium independently and a large part of the work, especially the isolation of the metallic thallium was done by Lamy, Crookes tried to secure his priority on the work. Lamy was awarded a medal at the International Exhibition in London 1862: For the discovery of a new and abundant source of thallium and after heavy protest Crookes also received medal: thallium, for the discovery of the new element. The controversy between both scientists continued through 1862 and 1863. Most of the discussion ended after Crookes was elected Fellow of the Royal Society in June 1863.[9] [10]
3
Thallium
4
Occurrence and production
Corroded thallium rod
Although the metal is reasonably abundant in the Earth's crust at a concentration estimated to be about 0.7 mg/kg, mostly in association with potassium minerals in clays, soils, and granites, it is not generally considered to be commercially recoverable from those forms. The major source of commercial thallium is the trace amounts found in copper, lead, zinc, and other sulfide ores.
Thallium is found in the minerals crookesite TlCu7Se4, hutchinsonite TlPbAs5S9, and lorandite TlAsS2. It also occurs as trace in pyrite and extracted as a by-product of roasting this ore for sulfuric acid production.[2] The metal can be obtained from the smelting of lead and zinc rich ores. Manganese nodules found on the ocean floor also contain thallium, but nodule extraction is prohibitively expensive and potentially environmentally destructive. In addition, several other thallium minerals, containing 16% to 60% thallium, occur in nature as sulfide or selenide complexes with antimony, arsenic, copper, lead, and silver, but are rare, and have no commercial importance as sources of this element. Thallium metal can also be obtained as a by-product in the production of sulfuric acid by roasting of pyrite.[2] [11]
Isotopes Thallium has 25 isotopes which have atomic masses that range from 184 to 210. 203Tl and 205 Tl are the only stable isotopes, and 204Tl is the most stable radioisotope, with a half-life of 3.78 years. 202
Tl (half life 12.23 days) can be made in a cyclotron,[12] while 204Tl (half life 3.78 years) is made by the neutron activation of stable thallium in a nuclear reactor.[13]
Compounds Fluorides: Thallium(I) fluoride (TlF), Thallium(III) fluoride (TlF3) Chlorides: Thallium(I) chloride (TlCl), Thallium(II) chloride (TlCl2), Thallium(III) chloride (TlCl3) Bromides: Thallium(I) bromide (TlBr), Thallium(II) bromide (Tl2Br4) Iodides: Thallium triiodide (TlI), Thallium triiodide (TlI3) Hydrides: none listed Oxides: Thallium oxide (Tl2O), Thallium(III) oxide (Tl2O3) Sulfides: Thallium(I) sulfide Tl2S Selenides: Thallium(I) selenide Tl2Se Tellurides: none listed Nitrides: none listed
Thallium
Applications The odorless and tasteless thallium sulfate was once widely used as rat poison and ant killer. Since 1975, this use in the United States and many other countries is prohibited due to safety concerns.[2] Other uses: • thallium(I) sulfide's electrical conductivity changes with exposure to infrared light therefore making this compound useful in photocells.[14] [15] • thallium(III) salts, as thallium trinitrate or triacetate, are useful reagents in organic synthesis performing different transformations in aromatics, ketones, olefins, among others. • Thallium(I) bromide and thallium(I) iodide crystals have been used as infrared optical materials, because they are harder than other common infrared optics, and because they have transmission at significantly longer wavelengths. The trade name KRS-5 refers to [14] this material. • used in semiconductor materials for selenium rectifiers,[14] • used as a dopant for sodium iodide crystals in gamma radiation detection equipment, such as scintillation counters, • high-density liquid used for sink-float separation of minerals, • used in the treatment of ringworm and other skin infections. However this use has been limited due to the narrow therapeutic index.[14] • radioactive thallium-201 (half-life of 73 hours) is used for diagnostic purposes in nuclear medicine, particularly in stress tests used for risk stratification in patients with coronary artery disease A(CAD).[16] [17] This isotope of thallium can be generated using a transportable generator which is similar to the technetium cow.[18] The generator contains lead-201 (half life 9.33 hours) which decays by electron capture to the thallium-201. The lead-201 can be produced in a cyclotron by the bombardment of thallium with protons or deuterons by the (p,3n) and (d,4n) reactions.[19] • Thallium oxide has been used to manufacture glasses that have a high index of refraction. Combined with sulfur or selenium and arsenic, thallium has been used in the production of high-density glasses that have low melting points in the range of 125 and 150 °C. These glasses have room temperature properties that are similar to ordinary glasses and are durable, insoluble in water and have unique refractive indices.[14] • A mercury-thallium alloy, which forms a eutectic at 8.5% thallium, is reported to freeze at –60 °C, some 20 °C below the freezing point of mercury. This alloy is used in thermometers and low-temperature switches.[14] • thallium is used in the electrodes in dissolved oxygen analyzers.[2] • thallium is a constituent of the alloy in the anode plates in magnesium seawater batteries.[2] In addition, research activity with thallium is ongoing to develop high-temperature superconducting materials for such applications as magnetic resonance imaging, storage of magnetic energy, magnetic propulsion, and electric power generation and transmission. After the discovery of the first thallium barium calcium copper oxide superconductor in 1988 the research in applications started.[20]
5
Thallium
6
Toxicity Thallium and its compounds are extremely toxic, and should be handled with great care. Contact with skin is dangerous, and adequate ventilation should be provided when melting this metal. Thallium(I) compounds have a high aqueous solubility and are readily absorbed through the skin. Exposure to them should not exceed 0.1 mg per m² of skin in an 8-hour time-weighted average (40-hour work week). Thallium is a suspected human carcinogen.[21] .
Treatment and internal decontamination One of the main methods of removing thallium (both radioactive and normal) from humans is to use Prussian blue, which is a solid ion exchange material which absorbs thallium and releases potassium. Up to 20 g per day of Prussian blue is fed by mouth to the person, and it passes through their digestive system and comes out in the stool. Hemodialysis and hemoperfusion are also used to remove thallium from the blood serum. At later stage of the treatment additional potassium is used to mobilize thallium from the tissue.[22] [23]
Bioconcentration According to the United States Environmental Protection Agency (EPA), thallium release to the environment was reported in Texas and Ohio. This may indicate bioconcentration in aquatic ecosystems.[24]
See also • Thallium poisoning • Thallium compounds
External links • • • • •
WebElements.com — Thallium [25] pure Thallium >=99,99% picture in the element collection from Heinrich Pniok Toxicity, Thallium [27] NLM Hazardous Substances Databank – Thallium, Elemental [28] ATSDR - ToxFAQs [29]*pure Thallium cast in acrylic for safe handling [30]
[26]
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] " Chemical fact sheet — Thallium (http:/ / www. speclab. com/ elements/ thallium. htm)". Spectrum Laboratories. April 2001. . Retrieved 2008-02-02. [3] Liddell & Scott, A Greek-English Lexicon (http:/ / perseus. mpiwg-berlin. mpg. de/ cgi-bin/ resolveform?doc=Perseus:text:1999. 04. 0057;layout=;query=toc;loc=qallo/ s), sub θαλλος (http:/ / perseus. mpiwg-berlin. mpg. de/ cgi-bin/ resolveform?lookup=qallos& type=begin& lang=greek& searchText=& options=Sort+ Results+ Alphabetically& . submit=Submit& formentry=1& lang=greek) [4] G. Kirchhoff, R. Bunsen (1861). "Chemische Analyse durch Spectralbeobachtungen". Annalen der Physik und Chemie 189 (7): 337–381. doi: 10.1002/andp.18611890702 (http:/ / dx. doi. org/ 10. 1002/ andp. 18611890702). [5] Crookes, William (1862 - 1863). " Preliminary Researches on Thallium (http:/ / www. jstor. org/ stable/ 112218)". Proceedings of the Royal Society of London, 12: 150–159. doi: 10.1098/rspl.1862.0030 (http:/ / dx.
Thallium doi. org/ 10. 1098/ rspl. 1862. 0030). . [6] Crookes, William (1863). " On Thallium (http:/ / www. jstor. org/ stable/ 108794)". Philosophical Transactions of the Royal Society of London, 153: 173–192. doi: 10.1098/rstl.1863.0009 (http:/ / dx. doi. org/ 10. 1098/ rstl. 1863. 0009). . [7] DeKosky, Robert K. (1973). " Spectroscopy and the Elements in the Late Nineteenth Century: The Work of Sir William Crookes (http:/ / www. jstor. org/ stable/ 4025503)". The British Journal for the History of Science 6 (4): 400–423. doi: 10.1017/S0007087400012553 (http:/ / dx. doi. org/ 10. 1017/ S0007087400012553). . [8] Lamy, Claude-Auguste (1862). " De l'existencè d'un nouveau métal, le thallium (http:/ / gallica2. bnf. fr/ ark:/ 12148/ bpt6k30115. image. r=Comptes+ Rendus+ Hebdomadaires. f1254. langFR)". Comptes Rendus: 1255–. . [9] James, Frank A. J. L. (1984). " Of 'Medals and Muddles' the Context of the Discovery of Thallium: William Crookes's Early (http:/ / www. jstor. org/ stable/ 531576)". Notes and Records of the Royal Society of London 39 (1): 65–90. doi: 10.1098/rsnr.1984.0005 (http:/ / dx. doi. org/ 10. 1098/ rsnr. 1984. 0005). . [10] Emsley, John (2006). " Thallium (http:/ / books. google. de/ books?id=BACSR7TXWhoC)". The Elements of Murder: A History of Poison. Oxford University Press. pp. 326–327. ISBN 9780192806000. . [11] Downs, Anthony John (1993). " Chemistry of Aluminium, Gallium, Indium, and Thallium (http:/ / books. google. com/ books?id=v-04Kn758yIC)". Springer. pp. 89 and 106. . [12] Thallium Research (http:/ / www. eh. doe. gov/ ohre/ roadmap/ histories/ 0472/ 0472d. html) from Department of Energy [13] Manual for reactor produced radioisotopes (http:/ / www-pub. iaea. org/ MTCD/ publications/ PDF/ te_1340_web. pdf) from the International Atomic Energy Agency [14] C. R. Hammond. The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304857. [15] Nayer, P. S, Hamilton, O. (1977). " Thallium selenide infrared detector (http:/ / adsabs. harvard. edu/ abs/ 1977ApOpt. . 16. 2942N)". Appl. Opt. 16: 2942. doi: 10.1364/AO.16.002942 (http:/ / dx. doi. org/ 10. 1364/ AO. 16. 002942). . [16] Thallium Test (http:/ / www. wramc. amedd. army. mil/ departments/ nuclear/ PatientInfo/ Thallium. htm) from Walter Reed Army Medical Center [17] Thallium Stress Test (http:/ / www. americanheart. org/ presenter. jhtml?identifier=4743) from the American Heart Association [18] M. C., Lagunas-Solar; Little, F. E.; Goodart, C. D. (1982). " Abstract An integrally shielded transportable generator system for thallium-201 production (http:/ / www. medscape. com/ medline/ abstract/ 7169272)". International Journal of Applied Radiation Isotopes 33 (12): 1439–1443. doi: 10.1016/0020-708X(82)90183-1 (http:/ / dx. doi. org/ 10. 1016/ 0020-708X(82)90183-1). Abstract. [19] Thallium-201 production (http:/ / www. med. harvard. edu/ JPNM/ physics/ isotopes/ Tl/ Tl201/ prod. html) from Harvard Medical School's Joint Program in Nuclear Medicine [20] Sheng, Z. Z.; Hermann A. M. (1988). "Bulk superconductivity at 120 K in the Tl–Ca/Ba–Cu–O system". Nature 332: 138–139. doi: 10.1038/332138a0 (http:/ / dx. doi. org/ 10. 1038/ 332138a0). [21] " Biology of Thallium (http:/ / www. webelements. com/ webelements/ elements/ text/ Tl/ biol. html)". webelemnts. . Retrieved 2008-11-11. [22] Prussian blue fact sheet (http:/ / www. bt. cdc. gov/ radiation/ prussianblue. asp) from the Centers for Disease Control and Prevention [23] Malbrain, Manu L. N. G.; Lambrecht, Guy L. Y.; Zandijk, Erik; Demedts, Paul A.; Neels, Hugo M.; Lambert, Willy; De Leenheer, André P.; Lins, Robert L.; Daelemans, Ronny; (1997). "= Treatment of Severe Thallium Intoxication". Clinical Toxicology 35 (1): 97–100. doi: 10.3109/15563659709001173 (http:/ / dx. doi. org/ 10. 3109/ 15563659709001173). [24] " Technical Factsheet on: Thallium (http:/ / www. epa. gov/ ogwdw/ dwh/ t-ioc/ thallium. html)". . Retrieved 2009-06-06. [25] http:/ / www. webelements. com/ webelements/ elements/ text/ Tl/ index. html [26] http:/ / www. pse-mendelejew. de/ bilder/ tl. jpg [27] http:/ / www. emedicine. com/ emerg/ topic926. htm [28] http:/ / toxnet. nlm. nih. gov/ cgi-bin/ sis/ search/ r?dbs+ hsdb:@term+ @na+ @rel+ thallium,+ elemental [29] http:/ / www. atsdr. cdc. gov/ tfacts54. html [30] http:/ / www. smart-elements. com/ ?arg=zoom& element=Tl& art=1967& newitems=& ref=Tl& seite=0& total=1& suche=#magnify
7
Article Sources and Contributors
Article Sources and Contributors Thallium Source: http://en.wikipedia.org/w/index.php?oldid=307208954 Contributors: 21655, Aarchiba, Achaemenes, Aecis, Ageo020, Ahkond, Ahoerstemeier, Alchemist-hp, Alex.tan, Alvis, Andre Engels, Andros 1337, Antandrus, Arabani, Arcadian, Archimerged, Arkuat, Astanhope, Astatine-210, Bayou Banjo, Belg4mit, Benjah-bmm27, Benjamin Mako Hill, Binky, Blah27, BlueEarth, Blueshirts, Borishal, Brockert, Brutaldeluxe, Bryan Derksen, CSWarren, CWii, CYD, Carnildo, Cburnett, Cecil, Cgingold, Ch'marr, Chasingsol, Cmdr Scolan, Conversion script, Da monster under your bed, DabMachine, Dajwilkinson, Darrien, Davemcarlson, David Haslam, David Latapie, Delirium, Doggie389, Donarreiskoffer, Dr Zak, DrBob, Dschwen, Duja, Dysprosia, Edgar181, El C, Element16, Eleveneleven, Emperor, Emperorbma, Enok Walker, Eric119, Erik Zachte, Everyking, Excirial, Fabiform, Feline1, Femto, FlyingPenguins, Fvasconcellos, GT5162, Gene Nygaard, Gilliam, Gnfnrf, GrahamHardy, Greatpatton, Greg Lindahl, Grendelkhan, Grumpyoldgeek, Hak-kâ-ngìn, HappyCamper, HazyM, Hede2000, Helge Skjeveland, Hippietrail, Hofoen, Huji, Icairns, Ideyal, Itub, J.delanoy, JALockhart, Jacj, Jaelanrodriguez, Jaraalbe, Jd027, Jennyvu96, Jimbimedia, Joanjoc, Johantheghost, John, John Nevard, Josh, Josh Parris, Jsjxyz, KRSESQ, Karlhahn, Kcjenner, Kelisi, Kipholbeck, Kostmo, Kurykh, Kwamikagami, Liam Skoda, Lindmere, Linmhall, Lizzie Harrison, Looxix, Lpnsm1, Lumos3, Madder, MapsMan, Marc Venot, Materialscientist, Mattfiller, Matthew0028, Mav, Mccready, Michael Snyder, MightyWarrior, Minesweeper, Mortdefides, Mr Minchin, Mygerardromance, NReitzel, NT17, Nabokov, Namibnat, Neparis, Nergaal, Neverquick, Niro5, Nuclearmedzors, Ojs, Oliverkroll, Omicronpersei8, Oobopshark, Ortolan88, Oskar Wallströmer, OwenBlacker, Oxymoron83, Panthro, Paraballo, Petri Krohn, PhilHibbs, PigFlu Oink, PlatinumX, Poccil, Polonium, Poolkris, Proofreader77, Quadro, RTC, Raghav273, Remember, Roberta F., Romanm, Rossami, SHCGRA Max, Saperaud, Sbmehta, Schneelocke, Sengkang, Seraphimblade, Sewebster, ShaunMacPherson, Sheitan, ShelfSkewed, Shoessss, Shorvath, Skatebiker, Sl, Smrgeog, Snezzy, Snori, Sokkor, SpLoT, Spitfire, Stemonitis, Stephen Hodge, Stephenpratt, SteveRamone, Stifynsemons, Stone, Svante, Tagishsimon, Tetracube, The Anome, TheMoog, Thingg, Tiddly Tom, Tim Starling, Timtrent, Tomchiukc, TutterMouse, UnicornTapestry, Uthbrian, VMoreira, Versus22, Vsmith, Warut, Watch37264, Welte, Wendy Wendy, Widefox, Wtmitchell, Yekrats, Yilloslime, Yyy, Zhang He, Zundark, 337 anonymous edits
Image Sources, Licenses and Contributors image:Tl-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Tl-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Paddy, Saperaud Image: Thallium-croprotated.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Thallium-croprotated.jpg License: unknown Contributors: User:Materialscientist Image:Thallium rod corroded.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Thallium_rod_corroded.jpg License: unknown Contributors: User:Dschwen Image:Skull and crossbones.svg Source: http://en.wikipedia.org/w/index.php?title=File:Skull_and_crossbones.svg License: Public Domain Contributors: Andux, Bayo, Coyau, D0ktorz, Derbeth, Franzenshof, Ies, J.delanoy, Karelj, MarianSigler, Silsor, Stepshep, The Evil IP address, W!B:, 6 anonymous edits
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
8
Lead
1
Lead thallium ← lead → bismuth Sn ↑ Pb ↓ Uuq Periodic Table Extended Periodic Table General Name, symbol, number Element category Group, Period, Block Appearance
Standard atomic weight Electron configuration Electrons per shell
lead, Pb, 82 Post-transition metals 14, 6, p bluish gray
−1
207.2(1) g·mol 14
[Xe] 4f
10
5d
2
2
6s 6p
2, 8, 18, 32, 18, 4 Physical properties
Phase
solid Density (near r.t.)
11.34 g·cm−3
Liquid density at m.p.
10.66 g·cm−3
Melting point
600.61 K (327.46 °C, 621.43 °F) Boiling point
2022 K (1749 °C, 3180 °F)
Heat of fusion
4.77 kJ·mol−1
Heat of vaporization
179.5 kJ·mol−1
Specific heat capacity
(25 °C) 26.650 J·mol−1·K−1
Lead
2
Vapor pressure P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
978
1088
1229
1412
1660
2027
Atomic properties Crystal structure Oxidation states
face centered cubic 4, 2, -4 (Amphoteric oxide)
Electronegativity Ionization energies (more)
2.33 (Pauling scale) 1st: 715.6 kJ·mol−1 2nd: 1450.5 kJ·mol−1 −1
3rd: 3081.5 kJ·mol Atomic radius Covalent radius Van der Waals radius
175 pm 146±5 pm 202 pm Miscellaneous
Magnetic ordering
diamagnetic
Electrical resistivity
(20 °C) 208 n Ω·m
Thermal conductivity
(300 K) 35.3 W·m
Thermal expansion
(25 °C) 28.9 µm·m
Speed of sound (thin rod)
(r.t.) (annealed) −1 1190 m·s
Young's modulus
16 GPa
Shear modulus
5.6 GPa Bulk modulus
46 GPa
Poisson ratio
0.44
Mohs hardness
1.5 Brinell hardness
CAS registry number
38.3 MPa 7439-92-1 Most-stable isotopes
−1
·K
−1
−1
·K
−1
Lead
3
Main article: Isotopes of lead iso 204
Pb
205
Pb
NA 1.4% syn
half-life >1.4×1017 y 1.53×107 y
206
24.1%
206
207
22.1%
207
208
52.4%
208
Pb Pb Pb
210
Pb
trace
DM
DE (MeV)
DP
Alpha
2.186
200
Epsilon
0.051
205
Alpha
3.792
206
Beta
0.064
210
Hg Tl
Pb is stable with 124 neutron Pb is stable with 125 neutron Pb is stable with 126 neutron
22.3 y
Hg Bi
References
Lead (pronounced /ˈlɛd/) is a main-group element with symbol Pb (Latin: plumbum) and atomic number 82. Lead is a soft, malleable poor metal, also considered to be one of the heavy metals. Lead has a bluish-white color when freshly cut, but tarnishes to a dull grayish color when exposed to air. It has a shiny chrome-silver luster when melted into a liquid. Lead is used in building construction, lead-acid batteries, bullets and shot, weights, and is part of solder, pewter, fusible alloys and radiation shields. Lead has the highest atomic number of all stable elements, although the next element, bismuth, has a half-life so long (longer than the estimated age of the universe) it can be considered stable. Like mercury, another heavy metal, lead is a potent neurotoxin that accumulates in soft tissues and bone over time. Lead poisoning was documented in ancient Rome, Greece, and China.
Characteristics Lead is bright and silvery when freshly cut but the surface rapidly tarnishes in air to produce the more commonly observed dull luster normally associated with lead. It is a dense, ductile, very soft, highly malleable, bluish-white metal that has poor electrical conductivity. This true metal is highly resistant to corrosion, and because of this property, it is used to contain corrosive liquids (e.g., sulfuric acid). Because lead is very malleable and resistant to corrosion it is extensively used in building construction, e.g., external coverings of roofing joints. Lead can be toughened by adding a small amount of antimony or other metals to it. It is a common misconception that lead has a zero Thomson effect. All lead, except 204Pb, is the end product of a complex radioactive decay. Lead is also poisonous, as are its compounds, and therefore is dangerous to human health and use as a food containment device is not recommended.
Lead
4
History Lead has been commonly used for thousands of years because it is widespread, easy to extract and easy to work with. It is highly malleable and ductile as well as easy to smelt. Metallic lead beads dating back to 6400 B.C. have been found in Çatalhöyük in modern-day Turkey.[1] In the early Bronze Age, lead was used with antimony and arsenic. Lead is mentioned in the Book of Exodus (15:10). In alchemy, lead was thought to be the oldest metal and was associated with the planet Saturn. Lead pipes that bear the insignia of Roman emperors are still in service and many Roman "pigs" (ingots) of lead figure in Derbyshire lead mining history and in the history of the industry in other English centres. The Romans also used lead in molten form to secure iron pins that held together large limestone blocks in certain monumental buildings. Lead's symbol Pb is an abbreviation of its Latin name plumbum for soft metals; originally it was plumbum nigrum (literally, "black plumbum"), where plumbum candidum (literally, "bright plumbum") was tin. The English words "plumbing", "plumber", "plumb", and "plumb-bob" also derive from this Latin root.
Occurrence Metallic lead does occur in nature, but it is rare. Lead is usually found in ore with zinc, silver and (most abundantly) copper, and is extracted together with these metals. The main lead mineral is galena (PbS), which contains 86.6% lead. Other common varieties are cerussite (PbCO3) and anglesite (PbSO4).
Ore processing Most ores contain less than 10% lead, and ores containing as little as 3% lead can be economically exploited. Ores are crushed and concentrated by froth flotation typically to 70% or more. Sulfide ores are roasted, producing primarily lead oxide and a mixture of sulfates and silicates of lead and other metals contained in the ore.[2] Lead oxide from the roasting process is reduced in a coke-fired blast furnace.[3] This converts most of the Lead ore lead to its metallic form. Three additional layers separate in the process and float to the top of the metallic lead. These are slag (silicates containing 1.5% lead), matte (sulfides containing 15% lead), and speiss (arsenides of iron and copper). These wastes contain concentrations of copper, zinc, cadmium, and bismuth that can be recovered economically, as can their content of unreduced lead.[2] Metallic lead that results from the roasting and blast furnace processes still contains significant contaminants of arsenic, antimony, bismuth, zinc, copper, silver, and gold. The melt is treated in a reverberatory furnace with air, steam, and sulfur, which oxidizes the contaminants except silver, gold, and bismuth. The oxidized contaminants are removed by drossing, where they float to the top and are skimmed off.[2] [4] Most lead ores contain significant concentrations of silver, resulting in the smelted metal also containing silver as a contaminant. Metallic silver as well as gold is removed and recovered economically by means of the Parkes process.[2] [4] [5]
Lead Desilvered lead is freed of bismuth according to the Betterton-Kroll process by treating it with metallic calcium and magnesium, which forms a bismuth dross that can be skimmed off.[2] [4] Very pure lead can be obtained by processing smelted lead electrolytically by means of the Betts process. The process uses anodes of impure lead and cathodes of pure lead in an electrolyte of silica fluoride.[2] [4]
Production and recycling Production and consumption of lead is increasing worldwide. Total annual production is about 8 million tonnes; about half is produced from recycled scrap. Top lead producing countries, as of 2008, are Australia, China, USA, Peru, Canada, Mexico, Sweden, Morocco, South Africa and North Korea.[4] Australia, China and the United States account for more than half of primary production.[6] • 2008 mine production: 3,886,000 tones • 2008 metal production: 8,725,000 tones • 2008 metal consumption: 8,706,000 tones[7] At current use rates, the supply of lead is estimated to run out in 42 years.[8] Environmental analyst, Lester Brown, however, has suggested lead could run out within 18 years based on an extrapolation of 2% growth per year.[9] This may need to be reviewed to take account of renewed interest in recycling, and rapid progress in fuel cell technology.
Isotopes Lead has many isotopes but 4 stable ones. The 4 stable isotopes are204Pb, 206Pb, 207Pb and 208 Pb with204Pb regarded as primordial Pb and 206, 207, 208 are formed from decay of U and Th. The one common radiogenic isotope, 202Pb, has a half-life of approximately 53,000 years.[10]
Health effects Lead is a poisonous metal that can damage nervous connections (especially in young children) and cause blood and brain disorders. Because of its low reactivity and solubility, lead poisoning usually only occurs in cases when the lead is dispersed, like when sanding lead based paint, or long term exposure in the case of pewter tableware. Long term exposure to lead or its salts (especially soluble salts or the strong oxidant PbO2) can cause nephropathy, and colic-like abdominal pains. The effects of lead are the same whether it enters the body through breathing or swallowing. Lead can affect almost every organ and system in the body. The main target for lead toxicity is the nervous system, both in adults and children. Long-term exposure of adults can result in decreased performance in some tests that measure functions of the nervous system. It may also cause weakness in fingers, wrists, or ankles. Lead exposure also causes small increases in blood pressure, particularly in middle-aged and older people and can cause anemia. Exposure to high lead levels can severely damage the brain and kidneys in adults or children and ultimately cause death. In pregnant women, high levels of exposure to lead may cause miscarriage. Chronic, high-level exposure in men can damage the organs responsible for sperm production. The concern about lead's role in cognitive deficits in children has brought about widespread reduction in its use (lead exposure has been linked to learning disabilities[11] ). Most cases
5
Lead of adult elevated blood lead levels are workplace-related.[12] High blood levels are associated with delayed puberty in girls.[13] Lead has been shown many times to permanently reduce the cognitive capacity of children at extremely low levels of exposure.[14] There appears to be no detectable lower limit, below which lead has no effect on cognition. During the 20th century, the use of lead in paint pigments was sharply reduced because of the danger of lead poisoning, especially to children.[15] [16] [17] By the mid-1980s, a significant shift in lead end-use patterns had taken place. Much of this shift was a result of the U.S. lead consumers' compliance with environmental regulations that significantly reduced or eliminated the use of lead in non-battery products, including gasoline, paints, solders, and water systems. Lead use is being further curtailed by the European Union's RoHS directive. Lead may still be found in harmful quantities in stoneware, vinyl (such as that used for tubing and the insulation of electrical cords), and brass manufactured in China. Between 2006 and 2007 many children's toys made in China were recalled, primarily due to lead in paint used to color the product. Older houses may still contain substantial amounts of lead paint. White lead paint has been withdrawn from sale in industrialized countries, but the yellow lead chromate is still in use; for example, Holland Colours Holcolan Yellow. Old paint should not be stripped by sanding, as this produces inhalable dust. Lead salts used in pottery glazes have on occasion caused poisoning, when acidic drinks, such as fruit juices, have leached lead ions out of the glaze.[18] It has been suggested that what was known as "Devon colic" arose from the use of lead-lined presses to extract apple juice in the manufacture of cider. Lead is considered to be particularly harmful for women's ability to reproduce. For that reason, many universities do not hand out lead-containing samples to women for instructional laboratory analyses. Lead(II) acetate (also known as sugar of lead) was used by the Roman Empire as a sweetener for wine, and some consider this to be the cause of the dementia that affected many of the Roman Emperors.[19] Lead as a soil contaminant is a widespread issue, since lead is present in natural deposits and may also enter soil through (leaded) gasoline leaks from underground storage tanks or through a wastestream of lead paint or lead grindings from certain industrial operations. Lead can also be found listed as a criteria pollutant in the United States Clean Air Act section 108. Lead that is emitted into the atmosphere can be inhaled, or it can be ingested after it settles out of the air. It is rapidly absorbed into the bloodstream and is believed to have adverse effects on the central nervous system, the cardiovascular system, kidneys, and the immune system.[20]
Biochemistry of lead poisoning In the human body, lead inhibits porphobilinogen synthase and ferrochelatase, preventing both porphobilinogen formation and the incorporation of iron into protoporphyrin IX, the final step in heme synthesis. This causes ineffective heme synthesis and subsequent microcytic anemia. At lower levels, it acts as a calcium analog, interfering with ion channels during nerve conduction. This is one of the mechanisms by which it interferes with cognition. Acute lead poisoning is treated using disodium calcium edetate: the calcium chelate of the disodium salt of ethylene-diamine-tetracetic acid (EDTA). This chelating agent has a greater affinity for lead than for calcium and so the lead chelate is formed by exchange. This is then excreted in the urine leaving behind harmless calcium.[21]
6
Lead
7
Leaching of lead from metal surfaces It is clear from the Pourbaix diagram below that lead is more likely to corrode in a citrate medium than it is in a non-complexing medium. The central part of the diagram shows that lead metal oxidizes more easily in the citrate medium than in normal water.
The Pourbaix diagram for lead in a non-complexing aqueous medium (eg perchloric acid / sodium hydroxide)
The Pourbaix diagram for lead in citric acid/citrate
In a Pourbaix diagram, the acidity is plotted on the x axis using the pH scale, while how oxidising/reducing nature of the system is plotted on the y axis in terms of volts relative to the standard hydrogen electrode. The diagram shows the form of the element which is most chemically stable at each point, it only comments on thermodynamics and it says nothing about the rate of change (kinetics).
Occupational exposure It is widely used in the production of batteries, metal products (solder and pipes), ammunition and devices to shield X-rays leading to its exposure to the people working in these industries. Use of lead in gasoline, paints and ceramic products, caulking, and pipe solder has been dramatically reduced in recent years because of health concerns. Ingestion of contaminated food and drinking water is the most common source of lead exposure in humans. Exposure can also occur via inadvertent ingestion of contaminated soil/dust or lead-based paint.
Testing Water contamination can be tested with commercially available kits. Analysis of lead in whole blood is the most common and accurate method of assessing lead exposure in human. Erythrocyte protoporphyrin (EP) tests can also be used to measure lead exposure, but are not as sensitive at low blood lead levels (<20 μg/dL). Lead in blood reflects recent exposure. Bone lead measurements are an indicator of cumulative exposure. While measurements of urinary lead levels and hair have been used to assess lead exposure, they are not reliable.
Lead
8
Chemistry Various oxidized forms of lead are easily reduced to the metal. An example is heating PbO with mild organic reducing agents such as glucose. A mixture of the oxide and the sulfide heated together without any reducing agent will also form the metal.[5] 2 PbO + PbS → 3 Pb + SO2 Metallic lead is attacked only superficially by air, forming a thin layer of oxide that protects it from further oxidation. The metal is not attacked by sulfuric or hydrochloric acids. It does, however, dissolve in nitric acid with the evolution of nitric oxide gas to form dissolved Pb(NO3)2. 3 Pb + 8 H+ + 8 NO−3 → 3 Pb2+ + 6 NO−3 + 2 NO + 4 H2O
When heated with nitrates of alkali metals, metallic lead oxidizes to form PbO (also known as litharge), leaving the corresponding alkali nitrite. PbO is representative of lead's +2 oxidation state. It is soluble in nitric and acetic acids, from which solutions it is possible to precipitate halide, sulfate, chromate, carbonate (PbCO3), and basic carbonate (Pb3(OH)2(CO3)2) salts of lead. The sulfide can also be precipitated from acetate solutions. These salts are all poorly soluble in water. Among the halides, the iodide is less soluble than the bromide, which, in turn, is less soluble than the chloride.[22] Lead(II) oxide is also soluble in alkali hydroxide solutions to form the corresponding plumbite salt.[5] PbO + 2 OH− + H2O → Pb(OH)2−4 Chlorination of plumbite solutions causes the formation of lead's +4 oxidation state. Pb(OH)2−4 + Cl2 → PbO2 + 2 Cl− + 2 H2O Lead dioxide is representative of the +4 oxidation state, and is a powerful oxidizing agent. The chloride of this oxidation state is formed only with difficulty and decomposes readily into lead(II) chloride and chlorine gas. The bromide and iodide of lead(IV) are not known to exist.[22] Lead dioxide dissolves in alkali hydroxide solutions to form the corresponding plumbates.[5] PbO2 + 2 OH− + 2 H2O → Pb(OH)2−6 Lead also has an oxide with mixed +2 and +4 oxidation states, red lead (Pb3O4), also known as minium. Lead readily forms an equimolar alloy with sodium metal that reacts with alkyl halides to form organometallic compounds of lead such as tetraethyl lead.[23]
Lead
9
Chloride complexes Lead(II) forms a series of complexes with chloride, the formation of which alters the corrosion chemistry of the lead. This will tend to limit the solubility of lead in saline media.
Diagram showing the forms of lead in chloride [24] media
Equilibrium constants for aqueous lead chloride complexes at 25 °C Pb2+ + Cl− → PbCl+
K1 = 12.59
PbCl+ + Cl− → PbCl2
K2 = 14.45
PbCl2 + Cl− → PbCl3−
K3 = 3.98 ×10−1
PbCl3− + Cl− → PbCl42−
K4 = 8.92 × 10−2
Phase diagrams of solubilities Lead(II) sulfate is poorly soluble, as can be seen in the following diagram showing addition of SO42− to a solution containing 0.1M of Pb2+. The pH of the solution is 4.5, as above that, Pb2+ concentration can never reach 0.1M due to the formation of Pb(OH)2. Observe that Pb2+ solubility drops 10,000 fold as SO42− reaches 0.1M.
Plot showing aqueous concentration of dissolved Pb2+ as a function of SO42−
Lead
10
Here it can be seen that the addition of chloride can lower the solubility of lead, however in chloride rich media (such as aqua regia) the lead can become soluble again as anionic chlorocomplexes.
Diagram showing the solubility of lead in chloride media. The lead concentrations are plotted as a function of the total chloride present.
Pourbaix diagram for lead in chloride (0.1 M) media
Applications Because of its high density and resistance against corrosion, lead is used for the ballast keel of sailboats. Its high weight-to-volume ratio allows it to counterbalance the heeling effect of wind on the sails while at the same time occupying a small volume and thus offering the least underwater resistance. For the same reason its is used in scuba diving weight belts to counteract the diver's natural buoyancy and that of his equipment. It does not have the weight-to-volume ratio of many heavy metals, but its low cost increases its use in these and other applications. • Lead is a major constituent of the lead-acid battery used extensively as a car battery.[25] • Lead is used as a coloring element in ceramic glazes, notably in the colors red and yellow.[26] • Lead is used to form glazing bars for stained glass or other multi-lit windows. The practice has become less common, not for danger but for stylistic reasons. • Lead is frequently used in polyvinyl chloride (PVC) plastic, which coats electrical cords.[27] [28]
Roman lead water pipes with taps
• Lead is used as projectiles for firearms and fishing sinkers because of its density, low cost compared to alternative products and ease of use due to relatively low melting point.[29] • Lead or "sheet-lead" is used as a sound deadening layer in such areas as wall, floor and ceiling design in sound studios where levels of airborne and mechanically produced sound are targeted for reduction or virtual elimination.[30] [31]
Lead
11
• Lead is used in some candles to treat the wick to ensure a longer, more even burn. Because of the dangers, European and North American manufacturers use more expensive alternatives such as zinc.[32] [33] • Lead is used as shielding from radiation, e.g. in x-ray rooms.[34]
Lead pipe in Roman baths
• Molten lead is used as a coolant, eg. for lead cooled fast reactors.[35] • Lead glass is composed of 12-28% lead oxide. It changes the optical characteristics of the glass and reduces the transmission of radiation.[36] • Lead is the traditional base metal of organ pipes, mixed with varying amounts of tin to control the tone of the pipe.[37] [38] • Lead is used as electrodes in the process of electrolysis. • Lead is used in solder for electronics, although this usage is being phased out by some countries to reduce the amount of environmentally unfriendly waste. • Lead is used in high voltage power cables as sheathing material to prevent water diffusion into insulation. • Lead is added to brass to reduce machine tool wear. • Some artists using oil-based paints continue to use lead carbonate white, citing its properties in comparison with the alternatives. • Lead, in the form of strips or "tape" is used for the customization of tennis rackets. Tennis rackets of the past sometimes had lead added to them by the manufacturer to increase weight.[39] • Lead has many uses in the construction industry, e.g. lead sheets are used as architectural metals in roofing material, cladding, flashings, gutters and gutter joints, and on roof parapets. Detailed lead moldings are used as decorative motifs used to fix lead sheet. • Lead is still widely used in statues and sculptures. • Tetra-ethyl lead is used as an anti-knock additive for aviation fuel in piston driven aircraft. • Lead-based semiconductors, such as lead telluride, lead selenide and lead antimonide are finding applications in photovoltaic (solar energy) cells and infrared detectors.[40] • Lead is often used to balance the wheels of a car; this use is being phased out in favor of other materials for environmental reasons.
Lead
Former applications • Lead pigments were used in lead paint for white as well as yellow, orange, and red. Most have been discontinued due of the dangers of lead poisoning. However, lead chromate is still in industrial use. Lead carbonate (white) is the traditional pigment for the priming medium for oil painting, but it has been largely displaced by the zinc and titanium oxide pigments. It was also quickly replaced in water-based painting mediums. • Lead carbonate white was used by the Japanese geisha and in the West for face-whitening make-up, which caused ill-health in the wearer. • Lead was the hot metal used in hot metal typesetting. • Lead was used for plumbing in Ancient Rome. • Lead was used as a preservative for food and drink in Ancient Rome. • Lead was used for joining cast iron water pipes and used as a material for small diameter water pipes until the early 1970s. • Tetraethyl lead was used in leaded fuels to reduce engine knocking; however, this is no longer common practice in the Western world due to its incompatibility with catalytic converters. • The EPA banned the use of lead gasoline for highway transportation, beginning January 1, 1996.[41] • Lead has been used to make "clubs" or bats more lethal by melting it into a hole drilled into the top. • Lead was used to make bullets for slings. • Lead was used for shotgun pellets in the US until about 1992 when it was outlawed (for waterfowl hunting only) and replaced by "non-toxic" shot, primarily steel pellets. • Lead was used as a component of toys. Due to toy safety regulations, this use has been stopped in the United States. • Lead was used in car body filler, which was used in many custom cars in the 1940s–60s. Hence the term Leadsled. • Lead is a superconductor at 7.2 K and IBM tried to make a Josephson effect computer out of lead-alloy.[42] • Lead was also used in pesticides before the 1950s, when fruit orchards were treated (ATSDR). • A lead cylinder attached to a long line was used by sailors for the vital navigational task of determining water depth by 'heaving the lead' at regular internals. A soft tallow insert at it's base allowed the nature of the sea bed to be determined, further aiding position finding. Contrary to popular belief, pencil "leads" have never been made from lead. The term comes from the Roman stylus, called the penicillus, which was made of lead.[43] When the pencil originated as a wrapped graphite writing tool, the particular type of graphite being used was named plumbago (lit. "act for lead"; "lead mockup").
12
Lead
13
See also • • • •
Adult Blood Lead Epidemiology and Surveillance Lead-Free Toys Act Medical geology Plumbosolvency
Further reading • Keisch, B., Feller, R. L., Levine, A. S., and Edwards, R. R.: "Dating and Authenticating Works of Art by Measurement of Natural Alpha Emitters". In: Science, 155, No. 3767, p. 1238–1242, 1967. • Keisch, B: "Dating Works of Art Through their Natural Radioactivity: Improvements and Applications". In: Science, 160, p. 413–415, 1968. • Keisch, B: "Discriminating Radioactivity Measurements of Lead: New Tool for Authentication". In: Curator, 11, No. 1., p. 41–52, 1968. • Jose S. Casas and Jose Sordo. Lead Chemistry, Analytical Aspects. Environmental Impacts and Health Effects, 2006, First Edition, ELSEVIER B.V • R.M. Harrison and D.P.H. Laxen, Lead Pollution Causes and Control, 1981, First Publish, British Library Cataloguing in Publication Data
External links • ATSDR Case Studies in Environmental Medicine: Lead Toxicity Health and Human Services • Lead [45] at the Open Directory Project
[44]
U.S. Department of
pnb:ہسیس
References [1] Dennis L. Heskel (1983). " A Model for the Adoption of Metallurgy in the Ancient Middle East (http:/ / www. jstor. org/ stable/ 2742674)". Current Anthropology 24 (3): 362–366. doi: 10.1086/203007 (http:/ / dx. doi. org/ 10. 1086/ 203007). . [2] Samans, Carl H. (1949). Engineering Metals and their Alloys. MacMillan. [3] " Primary Extraction of Lead Technical Notes (http:/ / www. ldaint. org/ technotes1. htm)". LDA International. . Retrieved 7 April 2007. [4] " Primary Lead Refining Technical Notes (http:/ / www. ldaint. org/ technotes2. htm)". LDA International. . Retrieved 7 April 2007. [5] Linus, Pauling (1947 ed.). General Chemistry. W.H. Freeman. [6] " Lead Information (http:/ / www. ldaint. org/ information. htm)". LDA International. . Retrieved 2007-09-05. [7] " Lead and Zinc Statistics (http:/ / www. ilzsg. org/ static/ statistics. aspx?from=1)". International Lead and Zinc Study Group. . Retrieved 2009-02-19. [8] "How Long Will it Last?". New Scientist 194 (2605): 38–39. May 26, 2007. ISSN 0262-4079 (http:/ / worldcat. org/ issn/ 0262-4079). [9] Brown, Lester (2006). Plan B 2.0: Rescuing a Planet Under Stress and a Civilization in Trouble. New York: W.W. Norton. p. 109. ISBN 0393328317. [10] Georges, Audi (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3–128. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [11] Howard, Hu (1991). "Knowledge of diagnosis and reproductive history among survivors of childhood plumbism". American Journal of Public Health 81 (8): 1070–1072. doi: 10.2105/AJPH.81.8.1070 (http:/ / dx. doi. org/ 10. 2105/ AJPH. 81. 8. 1070). PMID 1854006. [12] " NIOSH ABLES (http:/ / www. cdc. gov/ niosh/ topics/ ABLES/ ables-description. html)". United States National Institute for Occupational Safety and Health. . Retrieved 2007-10-04.
Lead [13] Schoeters, G.; et al. (2008). " Endocrine Disruptors and Abnormalities of Pubertal Development (http:/ / www. blackwell-synergy. com/ doi/ abs/ 10. 1111/ j. 1742-7843. 2007. 00180. x)". Basic & Clinical Pharmacology & Toxicology 102: 168–175. doi:10.1111/j.1742-7843.2007.00180.x (inactive 2009-04-22) . . [14] Needleman, H.L. (11 Jan 1990). " The long-term effects of exposure to low doses of lead in childhood. An 11-year follow-up report (http:/ / content. nejm. org/ cgi/ content/ abstract/ 322/ 2/ 83)". New England Journal of Medicine 322 (2): 83–88. PMID 2294437. . [15] " NSW Multicultural Health Communication Service (http:/ / www. health. nsw. gov. au/ health-public-affairs/ mhcs/ publications/ 4465. html)". NSW Health. . Retrieved 7 April 2007. [16] " Download: Lead paint: Cautionary note (http:/ / www. epa. qld. gov. au/ publications?id=1528)". Queensland Government. . Retrieved 7 April 2007. [17] " Lead Paint Information (http:/ / www. qld. mpa. org. au/ index. php/ content/ 33/ )". Master Painters, Australia. . Retrieved 7 April 2007. [18] " Government report on lead poisoning from ceramic glazes (http:/ / www. fda. gov/ bbs/ topics/ CONSUMER/ CON00081. html)". . Retrieved 2008-04-24. [19] " The Pernicious Allure of Lead (http:/ / www. nytimes. com/ 2007/ 08/ 21/ science/ 21angi. html?_r=1& oref=slogin)". New York Times. . [20] Bergeson, Lynn L. (2008). "The proposed lead NAAQS: Is consideration of cost in the clean air act's future?". Environmental Quality Management 18: 79. doi: 10.1002/tqem.20197 (http:/ / dx. doi. org/ 10. 1002/ tqem. 20197). [21] Laurence, D.R. (1966). Clinical Pharmacology(Third Edition). [22] Brady, James E. and Holum, John R. (1996). Descriptive Chemistry of the Elements. John Wiley and Sons. ISBN 0471135577. [23] Windholz, Martha (1976 isbn = 0911910263). Merck Index of Chemicals and Drugs, 9th ed., monograph 8393. Merck. [24] Ignasi Puigdomenech, Hydra/Medusa Chemical Equilibrium Database and Plotting Software (2004) KTH Royal Institute of Technology, freely downloadable software at (http:/ / www. kemi. kth. se/ medusa/ ) [25] Stellman, Jeanne Mager (1998). Encyclopaedia of Occupational Health and Safety (http:/ / books. google. de/ books?id=nDhpLa1rl44C& pg=PT644). International Labour Organization. pp. 81.2–81.4. ISBN 9789221098164. . [26] Leonard, Alvin R.; Lynch, Glenn (1958). " Dishware as a Possible Source for Lead Poisoning (http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?tool=pmcentrez& artid=1512529)". Calif. Med. 89 (6): 414–416. [27] Zweifel, Hans (2009). Plastics Additives Handbook (http:/ / books. google. de/ books?id=WbBH5QFXOhoC& pg=PT475). Hanser Verlag. pp. 438. ISBN 9783446408012. . [28] C. E. Wilkes; J. W. Summers; C. A. Daniels (eds.). With contributions by: Mark T. Berard .... (2005). PVC handbook (http:/ / books. google. de/ books?id=YUkJNI9QYsUC& pg=PA106). München: Hanser. p. 106. ISBN 9781569903797. . [29] Rooney, Corinne. " Contamination at Shooting Ranges (http:/ / www. lead. org. au/ fs/ shootingranges. pdf)" (PDF). The Lead Group, incorporated. . Retrieved 7 April 2007. [30] Sivaraman Guruswamy. (2000). Engineering properties and applications of lead alloys (http:/ / books. google. de/ books?id=gDwOAAAAQAAJ& pg=PA31). New York, NY: Marcel Dekker. p. 31. ISBN 9780824782474. . [31] edited by Richard Lansdown and William Yule. (1986). The Lead debate : the environment, toxicology, and child health (http:/ / books. google. de/ books?id=TtGmjOv9CUAC& pg=PA240). London: Croom Helm. p. 240. ISBN 9780709916536. . [32] Randerson, James (June 2002). " Candle pollution (http:/ / www. newscientist. com/ article/ mg17423481. 900-candle-pollution. html)". NewScientist.com (2348). . Retrieved 2007-04-07. [33] . doi: doi:10.1016/S0048-9697(00)00359-4 (http:/ / dx. doi. org/ doi:10. 1016/ S0048-9697(00)00359-4). [34] Structural shielding design for medical X-ray imaging facilities. (http:/ / books. google. de/ books?id=DKu4YDjEluoC& pg=PA16). Bethesda, MD: National Council on Radiation Protection and Measurement. 2004. p. 16–17. ISBN 9780929600833. . [35] Tuček, Kamil (2006). " Comparison of sodium and lead-cooled fast reactors regarding reactor physics aspects, severe safety and economical issues (http:/ / www. ecolo. org/ documents/ documents_in_english/ SFRvsLFR-05. pdf)". Nuclear Engineering and Design 236: 1589. doi: 10.1016/j.nucengdes.2006.04.019 (http:/ / dx. doi. org/ 10. 1016/ j. nucengdes. 2006. 04. 019). . [36] Amstock, Joseph S. (1997). Handbook of glass in construction. McGraw-Hill Professional. ISBN 9780070016194. [37] The Art of Organ Building, Vol. 2 (http:/ / books. google. de/ books?id=I0h525OVoTgC& pg=PA503). pp. 250–251. ISBN 9780486213156. . [38] The Organ (http:/ / books. google. de/ books?id=cgDJaeFFUPoC& pg=PA412). pp. 412–413. ISBN 9780415941747. .
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Lead [39] Routledge Handbook of Biomechanics and Human Movement Science (http:/ / books. google. de/ books?id=p7Ne5IjK5H0C& pg=PA250). p. 250. ISBN 9780415408813. . [40] " Applications for Lead (http:/ / www. americanelements. com/ pb. html)". . Retrieved 7 April 2007. [41] " Banning of Leaded Gasoline for Highway Use (http:/ / www. accessmylibrary. com/ coms2/ summary_0286-6346110_ITM)". . Retrieved 23 September 2008. [42] Henkels, W. H.; Geppert, L. M.; Kadlec, J.; Epperlein, P. W.; Beha, H. (September 1985). " Josephson 4 K-bit cache memory design for a prototype signal processor. (http:/ / adsabs. harvard. edu/ abs/ 1985JAP. . . . 58. 2371H)". Harvard University. . Retrieved 7 April 2007. [43] " A history of pencils (http:/ / www. pencils. com/ history. html)". www.pencils.com. . Retrieved 7 April 2007. [44] http:/ / www. atsdr. cdc. gov/ csem/ lead/ [45] http:/ / www. dmoz. org/ Science/ Chemistry/ Elements/ Lead/ /
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Image Sources, Licenses and Contributors image:Pb-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Pb-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Paddy, Saperaud Image: Lead brick.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Lead_brick.jpg License: Public Domain Contributors: User:Changlc File:LeadOreUSGOV.jpg Source: http://en.wikipedia.org/w/index.php?title=File:LeadOreUSGOV.jpg License: Public Domain Contributors: Dejvid, Ra'ike, Saperaud Image:Pb in water Pourbiax diagram.png Source: http://en.wikipedia.org/w/index.php?title=File:Pb_in_water_Pourbiax_diagram.png License: Public Domain Contributors: Cadmium Image:Pb in citrate media pourbiax diagram.png Source: http://en.wikipedia.org/w/index.php?title=File:Pb_in_citrate_media_pourbiax_diagram.png License: Public Domain Contributors: Cadmium File:Lead complexes in chloride media.png Source: http://en.wikipedia.org/w/index.php?title=File:Lead_complexes_in_chloride_media.png License: Public Domain Contributors: Cadmium Image:PbSO4 solubility graph.png Source: http://en.wikipedia.org/w/index.php?title=File:PbSO4_solubility_graph.png License: Public Domain Contributors: User:Karlhahn Image:Lead sulphate pourdaix diagram.png Source: http://en.wikipedia.org/w/index.php?title=File:Lead_sulphate_pourdaix_diagram.png License: Public Domain Contributors: Cadmium Image:PbCl2 solubility graph.png Source: http://en.wikipedia.org/w/index.php?title=File:PbCl2_solubility_graph.png License: Public Domain Contributors: User:Karlhahn Image:Lead chloride pourdiax diagram.png Source: http://en.wikipedia.org/w/index.php?title=File:Lead_chloride_pourdiax_diagram.png License: Public Domain Contributors: Cadmium File:DSC00125 - Tubi di piombo romani - Foto di G. Dall'Orto.jpg Source: http://en.wikipedia.org/w/index.php?title=File:DSC00125_-_Tubi_di_piombo_romani_-_Foto_di_G._Dall'Orto.jpg License: unknown Contributors: User:G.dallorto/Palermo, user:g.dallorto File:Lead pipe Bath.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Lead_pipe_Bath.jpg License: unknown Contributors: User:Zureks
License
16
License
Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
17
Bismuth
1
Bismuth lead ← bismuth → poloniumSb ↑ Bi ↓ Uup
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
83Bi Periodic table
Appearance lustrous silver
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
941
1041
1165
1325
1538
1835
Bismuth
2
Atomic properties Oxidation states ElectronegativityIonization energies (more) 2nd: 1610 kJ·mol−1 3rd: 2466 kJ·mol−1Atomic radiusCovalent radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusPoisson ratioMohs hardnessBrinell hardnessCAS registry number Most stable isotopes Main article: Isotopes of bismuth iso
N.A.
half-life
DM
DE (MeV)
DP
207
syn
31.55 y
ε, β+
2.399
207
208
syn
368,000 y
ε, β+
2.880
208
209
100%
(19 ± 2) ×1018y
α
syn
3.04 ×106y
IT
Bi Bi Bi
210m
Bi
Pb Pb
205
Tl
0.271
210
Bi
bismuth, Bi, 83 poor metals15, 6, p208.98040(1) g·mol−1 [Xe] 4f14 5d10 6s2 6p3 2, 8, 18, 32, 18, 5 (Image) solid 9.78 g·cm−3 10.05 g·cm−3 544.7 K,271.5 °C,520.7 °F −1 −1 1837 K,1564 °C,2847 °F 11.30 kJ·mol 151 kJ·mol (25 °C) 25.52 J·mol−1·K−13, 5 (mildly acidic oxide) 2.02 (Pauling scale) 1st: 703 kJ·mol−1 156 pm 148±4 pm rhombohedral diamagnetic (20 °C) 1.29x10^-6Ω·m (300 K) 7.97 W·m−1·K−1 (25 °C) 13.4 µm·m−1·K−1 (20 °C) 1790 m/s 32 GPa 12 GPa 31 GPa 0.33 2.25 94.2 MPa 7440-69-9 Bismuth (pronounced /ˈbɪzməθ/) is a chemical element that has the symbol Bi and atomic number 83. This trivalent poor metal chemically resembles arsenic and antimony. Bismuth is heavy and brittle; it has a silvery white color with a pink tinge due to the surface oxide. Bismuth is the most naturally diamagnetic of all metals, and only mercury has a lower thermal conductivity. It is generally considered to be the last naturally occurring stable, non-radioactive element on the periodic table, although it is actually slightly radioactive, with an extremely long half-life. Bismuth compounds are used in cosmetics, medicines, and in medical procedures. As the toxicity of lead has become more apparent in recent years, alloy uses for bismuth metal as a replacement for lead have become an increasing part of bismuth's commercial importance.
Bismuth
3
Characteristics Bismuth is a brittle metal with a white, silver-pink hue, often occurring in its native form with an iridescent oxide tarnish showing many refractive colors from yellow to blue. When combusted with oxygen, bismuth burns with a blue flame and its oxide forms yellow fumes.[1] Its toxicity is much lower than that of its neighbors in the periodic table such as lead, tin, tellurium, antimony, and polonium.
Bismuth crystal with an iridescent oxide surface
Although ununpentium is theoretically more diamagnetic, no other metal is verified to be more naturally diamagnetic than bismuth.[1] (Superdiamagnetism is a different physical phenomenon.) Of any metal, it has the second lowest thermal conductivity (after mercury) and the highest Hall coefficient. It has a high electrical resistance.[1] When deposited in sufficiently thin layers on a substrate, bismuth is a semiconductor, rather than a
poor metal.[2] Elemental bismuth is one of very few substances of which the liquid phase is denser than its solid phase (water being the best-known example). Bismuth expands 3.32% on solidification; therefore, it was long an important component of low-melting typesetting alloys, which needed to expand to fill printing molds.[1] Though virtually unseen in nature, high-purity bismuth can form distinctive hopper crystals. These colorful laboratory creations are typically sold to collectors. Bismuth is relatively nontoxic and has a low melting point just above 271 °C, so crystals may be grown using a household stove, although the resulting crystals will tend to be lower quality than lab-grown crystals.
Isotopes While bismuth was traditionally regarded as the element with the heaviest stable isotope, bismuth-209, it had long been suspected to be unstable on theoretical grounds. This was finally demonstrated in 2003 when researchers at the Institut d'Astrophysique Spatiale in Bismuth crystals Orsay, France, measured the alpha emission half-life of 209 19 [3] Bi to be 1.9 x 10 years, over a billion times longer than the current estimated age of the universe. Owing to its extraordinarily long half-life, for nearly all applications bismuth can be treated as if it is stable and non-radioactive. The radioactivity is of academic interest, however, because bismuth is one of few elements whose radioactivity was suspected, and indeed theoretically predicted, before being detected in the laboratory.
Bismuth
4
History Bismuth (New Latin bisemutum from German Wismuth, perhaps from weiße Masse, "white mass") was confused in early times with tin and lead because of its resemblance to those elements. Bismuth has been known since ancient times, and so no one person is credited with its discovery. Agricola, in De Natura Fossilium states that bismuth is a distinct metal in a family of metals including tin and lead in 1546 based on observation of the metals and their physical properties.[4] Claude François Geoffroy demonstrated in 1753 that this metal is distinct from lead and tin.[1] "Artificial bismuth" was commonly used in place of the actual metal. It was made by hammering tin into thin plates, and cementing them by a mixture of white tartar, saltpeter, and arsenic, stratified in a crucible over an open fire. Bismuth was also known to the Incas and used (along with the usual copper and tin) in a special bronze alloy for knives.[5]
Occurrence and production New York prices Time
Price (USD/lb.)
December 2000
$3.85–$4.15
November 2002
$2.70–$3.10
December 2003
$2.60–$2.90
June 2004
$3.65–$4.00
September 2005
$4.20–$4.60
September 2006
$4.50–$4.75
November 2006
$6.00–$6.50
December 2006
$7.30–$7.80
March 2007
$9.25–$9.75
April 2007
$10.50–$11.00
June 2007
$18.00–$19.00
November 2007
$13.50–$15.00
Bismuth output in 2005
In the Earth's crust, bismuth is about twice as abundant as gold. It is not usually economical to mine it as a primary product. Rather, it is usually produced as a byproduct of the processing of other metal ores, especially lead, tungsten (China), tin, copper, and also silver (indirectly) or other metallic elements.
The most important ores of bismuth are bismuthinite and bismite.[1] In 2005, China was the top producer of bismuth with at least 40% of the world share followed by Mexico and Peru, reports the British Geological Survey. Native bismuth is known from Australia, Bolivia, and China.
Bismuth
5
According to the USGS, world 2006 bismuth mine production was 5,700 tonnes, of which China produced 3,000 tonnes, Mexico 1,180 tonnes, Peru 950 tonnes, and the balance Canada, Kazakhstan and other nations. World 2006 bismuth refinery production was 12,000 tonnes, of which China produced 8,500 tonnes, Mexico 1,180 tonnes, Belgium 800 tonnes, Peru 600 tonnes, Japan 510 tonnes, and the balance Canada and other nations.[6] Bismite mineral
The difference between world bismuth mine production and refinery production reflects bismuth's status as a byproduct metal. Bismuth travels in crude lead bullion (which can contain up to 10% bismuth) through several stages of refining, until it is removed by the Kroll-Betterton process or the Betts process. The Kroll-Betterton process uses a pyrometallurgical separation from molten lead of calcium-magnesium-bismuth drosses containing associated metals (silver, gold, zinc, some lead, copper, tellurium, and arsenic), which are removed by various fluxes and treatments to give high-purity bismuth metal (over 99% Bi). The Betts process takes cast anodes of lead bullion and electrolyzes them in a lead fluosilicate-hydrofluosilicic acid electrolyte to yield a pure lead cathode and an anode slime containing bismuth. Bismuth will behave similarly with another of its major metals, copper. Thus world bismuth production from refineries is a more complete and reliable statistic. According to the Bismuth Advocate News[] , the price for bismuth metal from year-end 2000 to September 2005 was stuck in a range from $2.60 to $4.15 per lb., but after this period the price started rising rapidly as global bismuth demand as a lead replacement and other uses grew rapidly. New mines in Canada and Vietnam may relieve the shortages, but prices are likely to remain above their previous level for the foreseeable future. The Customer-Input price for bismuth is more oriented to the ultimate consumer; it started January 2008 at US$39.40 per kilogram ($17.90 per pound) in January 2008 and reached US$35.55 per kg (US$16.15 per lb.) in September 2008. [7]
Recycling While bismuth is most available today as a byproduct, its sustainability is more dependent on recycling. Bismuth is mostly a byproduct of lead smelting, along with silver, zinc, antimony, and other metals, and also of tungsten production, along with molybdenum and tin, and also of copper and tin production. Recycling bismuth is difficult in many of its end uses, primarily because of scattering. Probably the easiest to recycle would be bismuth-containing fusible alloys in the form of larger objects, then larger soldered objects. Half of the world solder consumption is in electronics (i.e., circuit boards).[8] As the soldered objects get smaller or contain little solder or little bismuth, the recovery gets progressively more difficult and less economic, although solder with a sizable silver content will be more worth recovering. Next in recycling feasibility would be sizeable catalysts with a fair bismuth content, perhaps as bismuth phosphomolybdate, and then bismuth used in galvanizing and as a free-machining metallurgical additive. Finally, the bismuth in the uses where it gets scattered the most, in stomach medicines (bismuth subsalicylate), paints (bismuth vanadate) on a dry surface, pearlescent cosmetics (bismuth oxychloride), and
Bismuth bismuth-containing bullets. The bismuth is so scattered in these uses as to be unrecoverable with present technology. Bismuth can also be available sustainably from greater efficiency of use or substitution, most likely stimulated by a rising price. For the stomach medicine, another active ingredient could be substituted for some or all of the bismuth compound . It would be more difficult to find an alternative to bismuth oxychloride in cosmetics to give the pearlescent effect. However, there are many alloying formulas for solders and therefore many alternatives. The most important sustainability fact about bismuth is its byproduct status, which can either improve sustainability (i.e., vanadium or manganese nodules) or, for bismuth from lead ore, constrain it; bismuth is constrained. The extent that the constraint on bismuth can be ameliorated or not is going to be tested by the future of the lead storage battery, since 90% of the world market for lead is in storage batteries for gasoline or diesel-powered motor vehicles. A huge worldwide research and development effort is underway to develop and manufacture advanced (non-lead) batteries or electrical storage devices. In December 2008, it was announced that a consortium of 14 U.S. firms were seeking $1 billion in Federal aid to build a battery manufacturing plant to make Li-on batteries for electric automobiles. The plant would build the basic cell and the individual consortium firm would finish its batteries by adding their own electronics, changing the voltage, chemistry, and other parameters as required. Similar batteries are already being manufactured in several countries in the Far East, and four dozen advanced battery plants are under construction in China. These batteries certainly foreshadow such batteries for gasoline-powered motor vehicles. When the Li-on battery or other electrical storage devices replace the lead storage battery, the bismuth recovered from lead ore is going to drop along with the production of lead ore and bismuth is going to become more dependent on recycling. The life-cycle assessment of bismuth will focus on solders, one of the major uses of bismuth, and the one with the most complete information. The average primary energy use for solders is around 200 MJ per kg, with the high-bismuth solder (58% Bi) only 20% of that value, and three low-bismuth solders (2% to 5% Bi) running very close to the average. The global warming potential averaged 10 to 14 kg carbon dioxide, with the high-bismuth solder about two-thirds of that and the low-bismuth solders about average. The acidification potential for the solders is around 0.9 to 1.1 kg sulfur dioxide equivalent, with the high-bismuth solder and one low-bismuth solder only one-tenth of the average and the other low-bismuth solders about average. [9] There is very little life-cycle information on other bismuth alloys or compounds.
Chemistry Bismuth forms trivalent and pentavalent compounds. The trivalent compounds are more common. Many of its chemical properties are similar to other elements in its group; namely, arsenic and antimony. Bismuth is stable to both dry and moist air at ordinary temperatures. At elevated temperatures, the vapours of the metal combine rapidly with oxygen, forming the yellow trioxide, Bi2O3.[10] On reaction with base, this oxide forms two series of oxyanions: BiO−2, which is polymeric and forms linear chains, and BiO3−3. The anion in Li3BiO3 is actually a cubic octameric anion, Bi8O24−24, whereas the anion in Na3BiO3 is tetrameric.[11] Bismuth sulfide, Bi2S3, occurs naturally in bismuth ores.[12] It is also produced by the combination of molten bismuth and sulfur.[10]
6
Bismuth
7
Unlike earlier members of group 15 elements such as nitrogen, phosphorus, and arsenic, and similar to the previous group 15 element antimony, bismuth does not form a stable hydride analogous to ammonia and phosphine. Bismuth hydride, bismuthine (BiH3), is an endothermic compound that spontaneously decomposes at room temperature. It is stable [11] only below −60°C. The halides of bismuth in low oxidation states have been shown to have unusual structures. What was originally thought to be bismuth(I) chloride, BiCl, turns out to be a complex compound consisting of Bi5+9 cations and BiCl2−5 and Bi2Cl2−8 anions.[11] [13] The Bi5+9 cation is also found in Bi10HfCl18, prepared by reducing a mixture of hafnium(IV) chloride and bismuth chloride with elemental bismuth. Other polyatomic bismuth cations are also known, such as Bi2+8, found in Bi8(AlCl4)2.[13] Bismuth also forms a low-valence bromide with the same structure as "BiCl". There is a true monoiodide, BiI, which contains chains of Bi4I4 units. BiI decomposes upon heating to the triiodide, BiI3, and elemental bismuth. A monobromide of the same structure also exists.[11] In oxidation state +3, bismuth forms trihalides with all of the halogens: BiF3, BiCl3, BiBr3, and BiI3. All of these, except BiF3, are hydrolysed by water to form the bismuthyl cation, BiO+, a commonly encountered bismuth oxycation.[11] Bismuth(III) chloride reacts with hydrogen chloride in ether solution to produce the acid HBiCl4.[14] Bismuth dissolves in nitric acid to form bismuth(III) nitrate, Bi(NO3)3. In the presence of excess water or the addition of a base, the Bi3+ ion reacts with the water to form BiO+, which precipitates as (BiO)NO3.[15] The oxidation state +5 is less frequently encountered. One such compound is the pentafluoride, BiF5, a powerful oxidising agent capable of oxidising xenon tetrafluoride:[14] BiF5 + XeF4 → XeF+3BiF−6 The dark red bismuth(V) oxide, Bi2O5, is unstable, liberating O2 gas upon heating.[16] In aqueous solution, the Bi3+ ion exists in various states of hydration, depending on the pH: pH range
Species
<3
Bi(H2O)3+6
0-4
Bi(H2O)5OH2+
1-5
Bi(H2O)4(OH)2+
5-14
Bi(H2O)3(OH)3
>11
Bi(H2O)2(OH)4−
These mononuclear species are in equilibrium. Polynuclear species also exist, the most important of which is BiO+, which exists in hexameric form as the octahedral complex [Bi6O4(OH)4]6+ (or 6 [BiO+]·2 H2O).[17]
Applications Bismuth oxychloride is sometimes used in cosmetics. Bismuth subnitrate and bismuth subcarbonate are used in medicine.[1] Bismuth subsalicylate (the active ingredient in Pepto-Bismol and (modern) Kaopectate) is used as an antidiarrheal and to treat some other gastro-intestinal diseases (oligodynamic effect). Also, the product Bibrocathol is an organic molecule containing Bismuth and is used to treat eye infections. Bismuth subgallate (the
Bismuth active ingredient in Devrom) is used as an internal deodorant to treat malodor from flatulence (or gas) and faeces. Some other current uses: • New research by physicists has found a potential role of bismuth as an ingredient in electronic circuits and in manufacturing next-generation solar cells which would have a greater efficiency. Bismuth allows for the creation of new diodes that can reverse their direction of current flow.[18] Ordinary semiconductor diodes are fixed in their direction. • Many bismuth alloys have low melting points and are widely used for fire detection and suppression system safety devices.[1] • Bismuth is used as an alloying agent in production of malleable irons.[1] • It is also used a thermocouple material.[1] • A carrier for U-235 or U-233 fuel in nuclear reactors.[1] • Bismuth has also been used in solders. The fact that bismuth and many of its alloys expand slightly when they solidify make them ideal for this purpose. • Bismuth subnitrate is a component of glazes that produces an iridescent luster finish. • Bismuth telluride is an excellent thermoelectric material; it is widely used. • A replacement propellant for xenon in Hall effect thrusters • In 1997 an antibody conjugate with Bi-213, which has a 45 minute half-life, and decays with the emission of an alpha-particle, was used to treat patients with leukemia. • In 2001, Professor Barry Allen and Dr. Graeme Melville at St. George Hospital in Sydney successfully produced Bi-213 in linac experiments which involved bombarding radium with bremsstrahlung photons. This cancer research team used Bi-213 in its Targeted Alpha Therapy (TAT) program. • The delta form of bismuth oxide when it exists at room temperature is a solid electrolyte for oxygen. This form normally only exists above and breaks down below a high temperature threshold, but can be electrodeposited well below this temperature in a highly alkaline solution. In the early 1990s, research began to evaluate bismuth as a nontoxic replacement for lead in various applications: • As noted above, bismuth has been used in solders; its low toxicity will be especially important for solders to be used in food processing equipment and copper water pipes. • A pigment in artists' oil and acrylic paint (Bismuth Vanadate) • Ingredient in free-machining brasses for plumbing applications • Ingredient in free-machining steels for precision machining properties • A catalyst for making acrylic fibers[1] • In low-melting alloys used in fire detection and extinguishing systems • Ingredient in lubricating greases • Dense material for fishing sinkers • In crackling microstars (dragon's eggs) in pyrotechnics, as the oxide, subcarbonate, or subnitrate • Replacement for lead in shot and bullets. The UK, U.S., and many other countries now prohibit the use of lead shot for the hunting of wetland birds, as many birds are prone to lead poisoning due to mistaken ingestion of lead (instead of small stones and grit) to aid digestion. Bismuth-tin alloy shot is one alternative that provides similar ballistic performance to lead. (Another less expensive but also more poorly performing alternative is "steel" shot, which is actually soft iron.)
8
Bismuth
9
• Bismuth core bullets are also starting to appear for use in indoor shooting ranges, where fine particles of lead from bullets impacting the backstop can be a chronic toxic inhalant problem. Owing to bismuth's crystalline nature, the bismuth bullets shatter into a non-toxic powder on impact, making recovery and recycling easy. The lack of malleability does, however, make bismuth unsuitable for use in expanding hunting bullets. • Fabrique Nationale de Herstal uses bismuth in the projectiles for its FN 303 less-lethal riot gun. According to the USGS, U.S. bismuth consumption in 2006 totaled 2,050 tonnes, of which chemicals (including pharmaceuticals, pigments, and cosmetics) were 510 tonnes, bismuth alloys 591 tonnes, metallurgical additives 923 tonnes, and the balance other uses.
Precautions Bismuth is not known to be toxic, compared to its periodic table neighbours (lead, antimony, and polonium), although some compounds (including bismuth chloride due to its corrosive acidity) are toxic and should be handled with care. As with lead, overexposure to bismuth can result in the formation of a black deposit on the gingiva, known as a bismuth line[19] . Fine bismuth powder can be pyrophoric.[20]
See also • Bismuth minerals
External links • WebElements.com - Bismuth [21] • USGS 2006 Minerals Yearbook: Bismuth • Bismuth Advocate News (BAN) [23]
[22]
• Bismuth Statistics and Information [24] - United States Geological Survey minerals information for bismuth • Laboratory growth of large crystals of Bismuth [25] by Jan Kihle Crystal Pulling Laboratories, Norway • Bismuth breaks half-life record for alpha decay [26] • Bismuth Crystals – Instructions & Pictures [27] pnb:ھتمسب
References [1] C. R. Hammond. The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304857. [2] C. A. Hoffman, J. R. Meyer, and F. J. Bartoli, A. Di Venere, X. J. Yi, C. L. Hou, H. C. Wang, J. B. Ketterson, and G. K. Wong (1993). "Semimetal-to-semiconductor transition in bismuth thin films". Phys. Rev. B 48: 11431. doi: 10.1103/PhysRevB.48.11431 (http:/ / dx. doi. org/ 10. 1103/ PhysRevB. 48. 11431). [3] Marcillac, Pierre de; Noël Coron, Gérard Dambier, Jacques Leblanc, and Jean-Pierre Moalic (April 2003). "Experimental detection of α-particles from the radioactive decay of natural bismuth". Nature 422: 876–878. doi: 10.1038/nature01541 (http:/ / dx. doi. org/ 10. 1038/ nature01541). PMID 12712201. [4] Agricola, Georgious (1546 (oring.); 1955(trans)). De Natura Fossilium. New York: Mineralogical Society of America. pp. 178. [5] " Bismuth Bronze from Machu Picchu, Peru (http:/ / adsabs. harvard. edu/ abs/ 1984Sci. . . 223. . 585G)". .
Bismuth [6] Carlin, Jr., James F.. " Commodity Report 2006: Bismuth (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ bismuth/ myb1-2006-bismu. pdf)" (PDF). United States Geological Survey. . Retrieved 2009-02-08. [7] " Customer input prices (http:/ / customer-inputprices. blogspot. com)". . Retrieved 2009-02-08. [8] Taylor, Harold A.. Bismuth. Financial Times Executive Commodity Reports. p. 17. ISBN 1 84083 326 2. [9] " IKP, Department of Life-Cycle Engineering (http:/ / leadfree. ipc. org/ files/ RoHS_15. pdf)". . Retrieved 2009-05-05. [10] Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0070494398 [11] S. M. Godfrey; C. A. McAuliffe; A. G. Mackie; R. G. Pritchard (1998). Nicholas C. Norman. ed. Chemistry of arsenic, antimony, and bismuth. Springer. pp. 67-84. ISBN 075140389X. [12] Ira Remsen (1886). An Introduction to the Study of Chemistry. Henry Holt and Company. p. 363. [13] R. J. Gillespie; J. Passmore (1975). H. J. Emeléus, A. G. Sharp. ed. Advances in Inorganic Chemistry and Radiochemistry. Academic Press. pp. 77-78. ISBN 0120236176. [14] Hitomi Suzuki; Yoshihiro Matano (2001). Organobismuth chemistry. Elsevier. p. 8. ISBN 0444205284. [15] Charles Adolphe Wurtz; William Houston Greene; H. F. Keller (1880). William Houston Greene. ed. Elements of modern chemistry (6th ed.). J.B. Lippincott. p. 351. [16] Thomas Scott; Mary Eagleson (1994). Concise encyclopedia chemistry. Walter de Gruyter. p. 136. ISBN 3110114518. [17] Arnold F. Holleman; Egon Wiberg (2001). Nils Wiberg. ed. Inorganic chemistry. Academic Press. p. 771. ISBN 0123526515. [18] http:/ / www. sciencedaily. com/ releases/ 2009/ 02/ 090219141530. htm [19] " bismuth line (http:/ / medical-dictionary. thefreedictionary. com/ bismuth+ line)". Farlex, Inc.. . Retrieved 8 February 2008. [20] Patnaik, Patnaik (2002). Handbook of Inorganic Chemical Compounds. McGraw-Hill Professional. ISBN 0070494398. [21] [22] [23] [24] [25] [26] [27]
http:/ / www. webelements. com/ webelements/ elements/ text/ Bi/ index. html http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ bismuth/ myb1-2006-bismu. pdf http:/ / www. basicsmines. com/ bismuth/ index. html http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ bismuth/ http:/ / en. wikipedia. org/ wiki/ Image:Bismuth-501g. jpg http:/ / physicsweb. org/ article/ news/ 7/ 4/ 16 http:/ / www. amazingrust. com/ Experiments/ how_to/ Bismuth_Crystals. html
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Article Sources and Contributors
Article Sources and Contributors Bismuth Source: http://en.wikipedia.org/w/index.php?oldid=308317024 Contributors: Acroterion, Ahoerstemeier, Akulo, Alansohn, Alchemist-hp, Alex.muller, Anclation, Andy Christ, Anescient, Animum, Antandrus, Anwar saadat, Archimerged, Arkuat, Army1987, Athaler, Baccyak4H, Bearian, Beetstra, Benbest, BlueEarth, Bmhtayl, Bobo192, BorgQueen, Brian0918, Bryan Derksen, Bucketsofg, CRKingston, CTho, CYD, Can't sleep, clown will eat me, Carnildo, Cflm001, ChemGardener, ClanCC, Cometstyles, Conversion script, Crakkpot, Creedlac, Crystalhound, Cyde, DMahalko, Dale101usa, DanMS, Darth Panda, David Latapie, Dcooper, Deglr6328, Dennisthe2, Dougher, Dprokopchuk, Dr Love Monkey, Dschwen, Eaolson, Eassin, Edgar181, Edisaloser, Egomaniac, Emperorbma, Enchanter, Epbr123, Epolk, Erik Zachte, Eszett, Excirial, Feedmecereal, Femto, Fluzwup, Furrykef, Fxer, Gdr, Gene Nygaard, Gene Ward Smith, Gilliam, Gioto, Gmel, Gnomishirish, Greatpatton, GregUbben, Grendelkhan, Gtstricky, Gwax, Gökhan, Hak-kâ-ngìn, Hamidbehbahani, He Who Lurks In Shadow, Herbee, HooXooH, Hvn0413, Iain99, Icairns, Ideyal, Intangir, Iridescent, JFlin5, JForget, JWB, Jagged 85, JanSöderback, Janahan, Jaraalbe, Jeendan, Joanjoc, JohnAldr, Johndburger, Jorgenumata, Jossi, JustAddPeter, Kas wiz, Keenan Pepper, Kelisi, Kingdon, Kjkolb, Koavf, Kurykh, Kwamikagami, Kwertii, LarryMorseDCOhio, Lars Washington, Lightmouse, LostLeviathan, Lradrama, LukeSurl, Luna Santin, MER-C, Madmarigold, Maitch, Marc Venot, Materialscientist, Mav, Mbralchenko, Melsaran, Mentifisto, Mervyn, Micha L. Rieser, Michal Nebyla, Mike Rosoft, Milkfish, Mincus, Minesweeper, Monazcordine, Montrealais, Muke, Nergaal, NickMartin, Nihiltres, No1lakersfan, Ocohen, Odie5533, PP Jewel, Pakaran, Paul R. Potts, Paul from Michigan, PaulHanson, Peacef5rog, Phoenix Hacker, Phoogenb, PierreAbbat, Plexust, Ponedonkey, Poolkris, Pras, Pwjb, Pyrochem, Qxz, R, R6144, RSido, RTC, Radiojon, Rallette, Reddi, Remember, Reneedownie, RexxS, Reza kalani, Rich Farmbrough, Roberta F., Romanm, Rominandreu, Rumping, Rursus, Sadads, Samf, Sandahl, Saperaud, Sbharris, Schneelocke, Sengkang, ShaunMacPherson, Shawn81, Shellreef, Sillybilly, Sl, Slicky, SpicyDragonZ, Spudtater, Squids and Chips, StevensonR, Stifynsemons, Stoic Squirrel, Stone, Supergirl54321, Tablizer, Tagishsimon, Tanaats, Terence, Terrx, Tetracube, TheNeutroniumAlchemist, Thricecube, Tim Starling, Tim.thelion, TitaniumDreads, Tj9991, V1adis1av, VASANTH S.N., Vsmith, Vuong Ngan Ha, WDavis1911, Walkerma, WarthogDemon, Willking1979, Wizard191, Wknight94, Yakbasser, Yekrats, Yonatan, Yoyoyo789, Yyy, Zizonus, 363 anonymous edits
Image Sources, Licenses and Contributors file:rhombohedral.svg Source: http://en.wikipedia.org/w/index.php?title=File:Rhombohedral.svg License: unknown Contributors: User:Stannered file:Electron shell 083 Bismuth.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_083_Bismuth.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 file:Bi chips.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Bi_chips.jpg License: unknown Contributors: User:Materialscientist File:Bi-crystal.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Bi-crystal.jpg License: unknown Contributors: User:Alchemist-hp, User:Richard Bartz Image:Bismuth-crystal.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Bismuth-crystal.jpg License: unknown Contributors: User:Micha L. Rieser Image:Bismuth (mined)2.PNG Source: http://en.wikipedia.org/w/index.php?title=File:Bismuth_(mined)2.PNG License: Public Domain Contributors: User:Anwar_saadat/bubble_maps_(FAQ) Image:Bismite.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Bismite.jpg License: unknown Contributors: Original uploader was Tillman at en.wikipedia
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11
Polonium
1
Polonium bismuth ← polonium → astatine Te ↑ Po ↓ Uuh Periodic Table Extended Periodic Table General Name, symbol, number
polonium, Po, 84
Element category
metalloids
Group, Period, Block
16, 6, p
Appearance
silvery Standard atomic weight
(209) g·mol
Electron configuration
[Xe] 6s 4f
−1
2
Electrons per shell
14
10
4
5d
6p
2, 8, 18, 32, 18, 6 Physical properties
Phase
solid Density (near r.t.)
(alpha) 9.196 g·cm
Density (near r.t.)
(beta) 9.398 g·cm
−3
−3
Melting point
527 K (254 °C, 489 °F) Boiling point
1235 K (962 °C, 1764 °F)
Heat of fusion
ca. 13 kJ·mol−1
Heat of vaporization
102.91 kJ·mol−1
Specific heat capacity
(25 °C) 26.4 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
at T(K)
100
1k
10 k
100 k
(846)
1003
1236
Atomic properties Crystal structure Oxidation states
cubic 6, 4, 2, -2 (amphoteric oxide)
Polonium
2 Electronegativity
Ionization energies
2.0 (Pauling scale) 1st: 812.1 kJ/mol
Atomic radius Covalent radius
168 pm 140±4 pm Miscellaneous
Magnetic ordering
nonmagnetic
Electrical resistivity
(0 °C) (α) 0.40 µΩ·m
Thermal conductivity
(300 K) ? 20 W·m−1·K−1
Thermal expansion
(25 °C) 23.5 µm·m−1·K−1
CAS registry number
7440-08-6 Most-stable isotopes
Main article: Isotopes of polonium iso 208
Po
NA syn
half-life 2.898 y
DM α
Po
syn
103 y
204
ε, β
1.401
208
α
4.979
205
ε, β
1.893
209
α
5.407
206
+
210
Po
syn
138.376 d
DP
5.215 +
209
DE (MeV)
Pb Bi Pb Bi Pb
References
Polonium (pronounced /pɵˈloʊniəm/) is a chemical element with the symbol Po and atomic number 84, discovered in 1898 by Marie and Pierre Curie. A rare and highly radioactive metalloid,[1] polonium is chemically similar to bismuth[2] and tellurium, and it occurs in uranium ores. Polonium has been studied for possible use in heating spacecraft. It is unstable; all isotopes of polonium are radioactive. Polonium is very volatile; it will almost completely vaporize at room temperatures.
Characteristics Isotopes Polonium has 27 known isotopes, all of which are radioactive. They have atomic masses that range from 194 to 218 u. 210Po (half-life 138.376 days) is the most widely available. 209Po (half-life 103 years) and 208Po (half-life 2.9 years) can be made through the alpha, proton, or deuteron bombardment of lead or bismuth in a cyclotron. 210
Po is an alpha emitter that has a half-life of 138.376 days; it decays directly to its stable daughter isotope, 206Pb. A milligram of 210Po emits about as many alpha particles per second as 4.5 grams of 226Ra. A few curies (1 curie equals 37 gigabecquerels, 1 Ci = 37 GBq) of 210Po emit a blue glow which is caused by excitation of surrounding air. A single gram of 210Po generates 140 watts of power.[3] Because it emits many alpha particles,
Polonium which are stopped within a very short distance in dense media and release their energy, 210 Po has been used as a lightweight heat source to power thermoelectric cells in artificial satellites; for instance, 210Po heat source was also used in each of the Lunokhod rovers deployed on the surface of the Moon, to keep their internal components warm during the [4] 210 lunar nights. Some anti-static brushes contain up to 500 microcuries (20 MBq) of Po as a source of charged particles for neutralizing static electricity in materials like photographic film.[5] 210 The majority of the time 210Po decays by emission of an alpha particle only, not by emission of an alpha particle and a gamma ray. About one in 100,000 alpha emissions causes an excitation in the nucleus which then results in the emission of a gamma ray.[6] This low gamma ray production rate (and the short range of alpha particles) makes it difficult to find and identify this isotope. Rather than gamma ray spectroscopy, alpha spectroscopy is the best method of measuring this isotope.
Solid state form Polonium is a radioactive element that exists in two metallic allotropes. The alpha form has a simple cubic crystal structure with an edge length of 335.2 picometres; the beta form is rhombohedral.[7] [8] The structure of polonoum has been characterized by X-ray diffraction [9] [10] and electron diffraction.[11]
Chemistry The chemistry of polonium is similar to that of tellurium The alpha form of solid polonium. and bismuth. Polonium dissolves readily in dilute acids, but is only slightly soluble in alkalis. The hydrogen compound PoH2 is liquid at room temperature (melting point -36.1°C, boiling point 35.3°C). Halides of the structure PoX2, PoX4 and PoX6 are known. The two oxides PoO2 and PoO3 are the products of oxidation of polonium.[12] 210
Po (in common with 238Pu) has the ability to become airborne with ease: if a sample is heated in air to 55 °C (131 °F), 50% of it is vaporized in 45 hours, even though the melting point of polonium is 254 °C (489 °F) and its boiling point is 962 °C (1763 °F).[13] More than one hypothesis exists for how polonium does this; one suggestion is that small clusters of polonium atoms are spalled off by the alpha decay. It has been reported that some microbes can methylate polonium by the action of methylcobalamin.[14] [15] This is similar to the way in which mercury, selenium and tellurium are methylated in living things to create organometallic compounds. As a result when considering the biochemistry of polonium one should consider the possibility that the polonium will follow the same biochemical pathways as selenium and tellurium.
3
Polonium
4
Compounds Oxides
Hydrides
Halogen Compounds
• PoO2 • PoO3
• PoH2
• PoX2, e.g. polonium dichloride, PoCl2 • PoX4 • PoX6
History Also tentatively called "Radium F", polonium was discovered by Marie Skłodowska-Curie and her husband Pierre Curie in 1898[16] and was later named after Marie Curie's native land of Poland (Latin: Polonia)[17] [18] Poland at the time was under Russian, Prussian, and Austrian partition, and did not exist as an independent country. It was Curie's hope that naming the element after her native land would publicize its lack of independence. Polonium may be the first element named to highlight a political controversy.[19] This element was the first one discovered by the Curies while they were investigating the cause of pitchblende radioactivity. The pitchblende, after removal of the radioactive elements uranium and thorium, was more radioactive than both the uranium and thorium put together. This spurred the Curies on to find additional radioactive elements. The Curies first separated out polonium from the pitchblende, and then within a few years, also isolated radium.
Detection Gamma counting By means of radiometric methods such as gamma spectroscopy (or a method using a chemical separation followed by an activity measurement with a non-energy-dispersive counter), it is possible to measure the concentrations of radioisotopes and to distinguish one from another. In practice, background Intensity against photon energy for noise would be present and depending on the detector, three isotopes. the line width would be larger which would make it harder to identify and measure the isotope. In biological/medical work it is common to use the natural 40K present in all tissues/body fluids as a check of the equipment and as an internal standard.
Polonium
5
Alpha counting The best way to test for (and measure) many alpha emitters is to use alpha-particle spectroscopy as it is common to place a drop of the test solution on a metal disk which is then dried out to give a uniform coating on the disk. This is then used as the test sample. If the thickness of the layer formed on the disk is too thick then the lines of the spectrum are broadened, this is because some of the energy of the alpha particles is lost during their movement through the layer of active material. An alternative method is to use internal liquid scintillation where the sample is mixed with a scintillation cocktail. When the light emitted is then counted, some machines will record the amount of light energy per radioactive decay event. Due to the imperfections of the liquid scintillation method (such as a failure for all the photons to be detected, cloudy or coloured samples can be difficult to count) and the fact that random quenching can reduce the number of photons generated per radioactive decay it is possible to get a broadening of the alpha spectra obtained through liquid scintillation. It is likely that these liquid scintillation spectra will be subject to a Gaussian broadening rather than the distortion exhibited when the layer of active material on a disk is too thick.
Intensity against alpha energy for four isotopes, note that the line width is narrow and the fine details can be seen.
Intensity against alpha energy for four isotopes, note that the line width is wide and some of the fine details can not be seen. This is for liquid scintillation counting where random effects cause a variation in the number of visible photons generated per alpha decay.
A third energy dispersive method for counting alpha particles is to use a semiconductor detector. From left to right the peaks are due to 209Po, 210Po, 239Pu and 241Am. The fact that isotopes such as 239Pu and 241Am have more than one alpha line indicates that the nucleus has the ability to be in different discrete energy levels (like a molecule can).
Occurrence and production Polonium is a very rare element in nature because of the short half-life of all its isotopes. It is found in uranium ores at about 100 micrograms per metric ton (1 part in 1010), which is approximately 0.2% of the abundance of radium. The amounts in the Earth's crust are not harmful. Polonium has been found in tobacco smoke from tobacco leaves grown with phosphate fertilizers.[20] [21] [22]
Polonium
Neutron capture Synthesis by (n,γ) reaction In 1934 an experiment showed that when natural 209Bi is bombarded with neutrons, 210Bi is created, which then decays to 210Po via β decay. The final purification is done pyrochemically followed by liquid-liquid extraction techniques.[23] Polonium may now be made in milligram amounts in this procedure which uses high neutron fluxes found in nuclear reactors. Only about 100 grams are produced each year, practically all of it in Russia, making polonium exceedingly rare.[24] [25]
Proton capture Synthesis by (p,n) and (p,2n) reactions It has been found that the longer-lived isotopes of polonium can be formed by proton bombardment of bismuth using a cyclotron. Other more neutron rich isotopes can be formed by the irradiation of platinum with carbon nuclei.[26]
Applications When it is mixed or alloyed with beryllium, polonium can be a neutron source: beryllium releases a neutron upon absorption of an alpha particle that is supplied by 210Po. It has been used in this capacity as a neutron trigger or initiator for nuclear weapons. However, a license is needed to own and operate this form of neutron source.[27] Other uses include the following. • Devices that eliminate static charges in textile mills and other places.[28] However, beta particle sources are more commonly used and are less dangerous. A non-radioactive alternative is to use a high-voltage DC power supply to ionise air positively or negatively as required.[29] • 210Po can be used as an atomic heat source to power radioisotope thermoelectric generators via thermoelectric materials.[30] • Because of its very high toxicity, polonium can be used as a poison (see, for example, Alexander Litvinenko poisoning). • Polonium is also used to eliminate dust on film.[31]
Toxicity Overview By mass, polonium-210 is around 250,000 times more toxic than hydrogen cyanide (the actual LD50 for 210Po is about 1 microgram for an 80 kg person (see below) compared with about 250 milligrams for hydrogen cyanide[32] ). The main hazard is its intense radioactivity (as an alpha emitter), which makes it very difficult to handle safely: one gram of Po will self-heat to a temperature of around 500 °C (932 °F).[3] Even in microgram amounts, handling 210Po is extremely dangerous, requiring specialized equipment and strict handling procedures. Alpha particles emitted by polonium will damage organic tissue easily if polonium is ingested, inhaled, or absorbed, although they do not penetrate the epidermis and hence are not hazardous if the polonium is outside the body.
6
Polonium
Acute effects The median lethal dose (LD50) for acute radiation exposure is generally about 4.5 Sv.[33] The committed effective dose equivalent 210Po is 0.51 µSv/Bq if ingested, and 2.5 µSv/Bq if inhaled.[34] Since 210Po has an activity of 166 TBq (4486.5 Ci) per gram[34] (1 gram produces 166×1012 decays per second), a fatal 4.5 Sv (J/kg) dose can be caused by ingesting 8.8 MBq (238 microcuries, µCi), about 50 nanograms (ng), or inhaling 1.8 MBq (48 µCi), about 10 ng. One gram of 210Po could thus in theory poison 20 million people of whom 10 million would die. The actual toxicity of 210Po is lower than these estimates, because radiation exposure that is spread out over several weeks (the biological half-life of polonium in humans is 30 to 50 days[35] ) is somewhat less damaging than an instantaneous dose. It has been estimated that a median lethal dose of 210Po is 0.015 GBq (0.4 mCi), or 0.089 micrograms, still an extremely small amount. [36] [37]
Long term (chronic) effects In addition to the acute effects, radiation exposure (both internal and external) carries a long-term risk of death from cancer of 5–10% per Sv.[33] The general population is exposed to small amounts of polonium as a radon daughter in indoor air; the isotopes 214Po and 218 Po are thought to cause the majority[38] of the estimated 15,000-22,000 lung cancer deaths in the US every year that have been attributed to indoor radon.[39] Tobacco smoking causes additional exposure to polonium.[40]
Regulatory exposure limits The maximum allowable body burden for ingested 210Po is only 1.1 kBq (30 nCi), which is equivalent to a particle massing only 6.8 picograms. The maximum permissible workplace concentration of airborne 210Po is about 10 Bq/m3 (3 × 10−10 µCi/cm³).[41] The target organs for polonium in humans are the spleen and liver.[42] As the spleen (150 g) and the liver (1.3 to 3 kg) are much smaller than the rest of the body, if the polonium is concentrated in these vital organs, it is a greater threat to life than the dose which would be suffered (on average) by the whole body if it were spread evenly throughout the body, in the same way as caesium or tritium (as T2O). 210
Po is widely used in industry, and readily available with little regulation or restriction. In the US, a tracking system run by the Nuclear Regulatory Commission will be implemented in 2007 to register purchases of more than 16 curies (590 GBq) of polonium 210 (enough to make up 5,000 lethal doses). The IAEA "is said to be considering tighter regulations... There is talk that it might tighten the polonium reporting requirement by a factor of 10, to 1.6 curies (59 GBq)."[43]
Famous poisoning cases Notably, the murder of Alexander Litvinenko, a Russian dissident, in 2006 was announced as due to 210Po poisoning[44] [45] (see Alexander Litvinenko poisoning). According to Prof. Nick Priest of Middlesex University, an environmental toxicologist and radiation expert, speaking on Sky News on December 2, Litvinenko was probably the first person ever to die of the acute α-radiation effects of 210Po.[46] The Polonium Restaurant (a Polish restaurant in Sheffield, England, owned by Boguslaw Sidorowicz and named after his folk band in the late 1970s) experienced increased interest and business as a result of internet searches for the phrase polonium restaurant.[47] [48] [49]
7
Polonium It has also been suggested that Irène Joliot-Curie was the first person ever to die from the radiation effects of polonium (due to a single intake) in 1956.[50] She was accidentally exposed to polonium in 1946 when a sealed capsule of the element exploded on her laboratory bench. A decade later, on 17 March 1956, she died in Paris from leukemia which may have been caused by that exposure. According to the book The Bomb in the Basement, several death cases in Israel during 1957-1969 were caused by 210Po.[51] A leak was discovered at a Weizmann Institute laboratory in 1957. Traces of 210Po were found on the hands of professor Dror Sadeh, a physicist who researched radioactive materials. Medical tests indicated no harm, but the tests did not include bone marrow. Sadeh died from cancer. One of his students died of leukemia, and two colleagues died after a few years, both from cancer. The issue was investigated secretly, and there was never any formal admission that a connection between the leak and the deaths had existed.[52]
Treatment It has been suggested that chelation agents such as British Anti-Lewisite (dimercaprol) can be used to decontaminate humans.[53] In one experiment, rats were given a fatal dose of 1.45 MBq/kg (8.7 ng/kg) of 210Po; all untreated rats were dead after 44 days, but 90% of the rats treated with the chelation agent HOEtTTC remained alive after 5 months.[54]
Commercial products containing polonium No nuclear authority has asserted that a commercial product was a likely source for the poisoning of Litvinenko. However, as Prof. Peter D. Zimmerman says, "Polonium 210 is surprisingly common. ...Polonium sources with about 10 percent of a lethal dose are readily available—even in a product sold on Amazon.com."[55] Potentially lethal amounts of polonium are present in anti-static brushes sold to photographers.[56] Many of the devices are available by mail order. General Electric markets a static eliminator module with 500 µCi (20 MBq), roughly 2.5 times the lethal dose of 210Po if 100%-ingested, for US$71;[57] Staticmaster sells replacement units with the same amount (500 µCi) of 210Po for US$36.[58] In USA, the devices with no more than 500 µCi of (sealed) 210Po per unit can be bought in any amount under a "general license"[59] which means that a buyer need not be registered by any authorities: the general license "is effective without the filing of an application with the Commission or the issuance of a licensing document to a particular person." If these sources were used to collect the amount of polonium likely used in the poisoning—and one could devise a method of separating the polonium from its protective casing—it would take 10–100 modules for price of US$360 to US$7,100. That such a thing could be done is extremely difficult according to the manufacturers and would be highly dangerous to anyone attempting to do so without some special equipment like a glovebox. Sometimes sources of polonium used in industry are stolen or lost. According to the Nuclear Regulatory Commission (NRC), there were registered at least 8 cases of loss of control of potentially lethal polonium sources in the USA during 2006. Tiny amounts of such radioisotopes are sometimes used in the laboratory and for teaching purposes—typically of the order of 4–40 kBq (0.1–1.0 µCi), in the form of sealed sources, with the polonium deposited on a substrate or in a resin or polymer matrix—are often
8
Polonium
9
exempt from licensing by the NRC and similar authorities as they are not considered hazardous. Small amounts of 210Po are available to the public in the United States by mail order from a company called United Nuclear as 'needle sources' for laboratory experimentation. It would require about 15,000 of these 210Po sources, at a total cost of about $1 million, to obtain a toxic quantity of polonium. They typically sell between four and eight sources per year.[60] [61] According to some estimates,[62] the cost of the quantity of pure polonium-210 used to kill Litvinenko would be around £20 million (US$39 million).[63] However, this estimation is based on retail prices of commercially available demonstration radiation sources with very small activities and cannot be considered as reasonable.
Tobacco [64] [65]
The presence of polonium in tobacco smoke has been known since the early 1960s. Some of the world's biggest tobacco firms researched ways to remove the substance—to no avail—over a 40-year period but never published the results.[22]
Radioactive polonium-210 contained in phosphate fertilizers is absorbed by the roots of plants (such as tobacco) and stored in its tissues.[66] [67] [68] Tobacco plants fertilized by rock phosphates contain polonium-210, which emits alpha radiation estimated to cause about 11,700 lung cancer deaths annually worldwide.[22] [69] [70]
See also • Decay chain • Polonium halo • Copernicium
External links • Build a pocket-sized ion chamber, useful for detecting Polonium
[71]
pnb:مینولوپ
References [1] " Chemical Elements.com - Metalloids (http:/ / www. chemicalelements. com/ groups/ metalloids. html)". . Retrieved 2009-05-05. [2] " Polonium (http:/ / hyperphysics. phy-astr. gsu. edu/ hbase/ pertab/ Po. html)". . Retrieved 2009-05-05. [3] " Polonium (http:/ / www. ead. anl. gov/ pub/ doc/ polonium. pdf)". Argonne National Laboratory. . Retrieved 2009-05-05. [4] Andrew Wilson (1987). Solar System Log. London: Jane's Publishing Company Ltd. p. 64. [5] " Staticmaster Ionizing Brushes (http:/ / www. amstat. com/ solutions/ staticmaster. html)". AMSTAT Industries. . Retrieved 2009-05-05. [6] " 210PO a decay (http:/ / atom. kaeri. re. kr/ cgi-bin/ decay?Po-210 A)". . Retrieved 2009-05-05. [7] Gary L. Miessler; Donald A. Tarr (2004). Inorganic Chemistry (3 ed.). Upper Saddle River, N.J.: Pearson Prentice Hall. pp. 285. ISBN 0-13-120198-0. [8] " The beta Po (A_i) Structure (http:/ / cst-www. nrl. navy. mil/ lattice/ struk/ a_i. html)". . Retrieved 2009-05-05. [9] Desando, R. J.; Lange, R. C. (1966). "The structures of polonium and its compounds—I α and β polonium metal". Journal of Inorganic and Nuclear Chemistry 28: 1837. doi: 10.1016/0022-1902(66)80270-1 (http:/ / dx. doi. org/ 10. 1016/ 0022-1902(66)80270-1). [10] Beamer, W. H.; Maxwell, C. R. (1946). "The Crystal Structure of Polonium". Journal of Chemical Physics 14: 569. doi: 10.1063/1.1724201 (http:/ / dx. doi. org/ 10. 1063/ 1. 1724201). [11] Rollier, M. A.; Hendricks, S. B.; Maxwell, L. R. (1936). "The Crystal Structure of Polonium by Electron Diffraction". Journal of Chemical Physics 4: 648. doi: 10.1063/1.1749762 (http:/ / dx. doi. org/ 10. 1063/ 1.
Polonium 1749762). [12] Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 0-12-352651-5. [13] Bogdan Wąs, Ryszard Misiak, Mirosław Bartyzel, Barbara Petelenz (2006). " Thermochromatographic Separation of 206,208Po from a Bismuth Target Bombardet with Protons (http:/ / www. ichtj. waw. pl/ ichtj/ nukleon/ back/ full/ vol51_2006/ v51s2p03f. pdf)". Nukleonica 51 (Suppl. 2): s3-s5. . [14] Momoshima N., Song L.X., Osaki S.,Maeda Y., (2001). " Formation and emission of volatile polonium compound by microbial activity and polonium methylation with methylcobalamin (http:/ / www. medscape. com/ medline/ abstract/ 11478248?prt=true)". Environ Sci Technol 35 (15): 2956–2960. doi: 10.1021/es001730+ (http:/ / dx. doi. org/ 10. 1021/ es001730+ ). . [15] Momoshima N., Song L.X., Osaki S.,Maeda Y., (2002). "Biologically induced Po emission from fresh water". J Environ Radioact. 63: 187. doi: 10.1016/S0265-931X(02)00028-0 (http:/ / dx. doi. org/ 10. 1016/ S0265-931X(02)00028-0). [16] Curie P., Curie M. (1898). "Radiations from Compounds of Uranium and of Thorium". Comptes Rendus 126: 1101. [17] Pfützner M. (1999). " Borders of the Nuclear World --- 100 Years After Discovery of Polonium (http:/ / adsabs. harvard. edu/ abs/ 1999AcPPB. . 30. 1197P)". Acta Physica Polonica B 30: 1197. . [18] Adloff J. P. (2003). "The centennial of the 1903 Nobel Prize for physics". Radichimica Acta 91: 681. doi: 10.1524/ract.91.12.681.23428 (http:/ / dx. doi. org/ 10. 1524/ ract. 91. 12. 681. 23428). [19] Kabzinska K. (1998). "Chemical and Polish aspects of polonium and radium discovery". Przemysl Chemiczny 77 (3): 104–107. [20] Kilthau, Gustave F. "Cancer risk in relation to radioactivity in tobacco". Radiologic Technology 67: 217–222. [21] " Alpha Radioactivity (210 Polonium) and Tobacco Smoke (http:/ / kidslink. bo. cnr. it/ besta/ fumo/ epolonio. html)". . Retrieved 2009-05-05. [22] Monique E. Muggli et al. (2008). "Waking a Sleeping Giant: The Tobacco Industry’s Response to the Polonium-210 Issue". American Journal of Public Health 98: 1643. doi: 10.2105/AJPH.2007.130963 (http:/ / dx. doi. org/ 10. 2105/ AJPH. 2007. 130963). [23] Schulz, Wallace W.; Schiefelbein, Gary F.; Bruns, Lester E. (1969). "Pyrochemical Extraction of Polonium from Irradiated Bismuth Metal". Ind. Eng. Chem. Process Des. Dev. 8 (4): 508. doi: 10.1021/i260032a013 (http:/ / dx. doi. org/ 10. 1021/ i260032a013). [24] " Q&A: Polonium-210 (http:/ / www. rsc. org/ chemistryworld/ News/ 2006/ November/ 27110601. asp)". RSC Chemistry World. 2006-11-27. . Retrieved 2009-01-12. [25] " Most Polonium Made Near the Volga River (http:/ / www. sptimesrussia. com/ index. php?action_id=2& story_id=20100)". The St. Petersburg Times - News. . [26] Atterling, H., Forsling, W. (1959). " Light Polonium Isotopes from Carbon Ion Bombardments of Platinum (http:/ / www. osti. gov/ energycitations/ product. biblio. jsp?osti_id=4238755)". Arkiv for Fysik 15 (1): 81–88. . [27] Rhodes, Richard (2002). Dark Sun: The Making of the Hydrogen Bomb. New York: Walker & Company. pp. 187–188. [28] " BBC News : College breaches radioactive regulations (http:/ / news. bbc. co. uk/ 1/ hi/ england/ 1868414. stm)". . Retrieved 2009-05-05. [29] " Static Control for Electronic Balance Systems (http:/ / www. thermo. com/ eThermo/ CMA/ PDFs/ Articles/ articlesFile_16929. pdf)". . Retrieved 2009-05-05. [30] Hanslmeier, Arnold (2002). The sun and space weather (http:/ / books. google. com/ books?id=07TEK_w3A4AC& pg=PA183). Springer. p. 183. ISBN 1402006845. . [31] Emsley, John (2001). Nature's Building Blocks. New York: Oxford University Press. p. 331. [32] " Hydrogen cyanide msds (http:/ / www. physchem. ox. ac. uk/ MSDS/ HY/ hydrogen_cyanide. html)". . [33] " Health Impacts from Acute Radiation Exposure (http:/ / www. pnl. gov/ main/ publications/ external/ technical_reports/ PNNL-14424. pdf)". . Retrieved 2009-05-05. [34] " Nuclide Safety Data Sheet: Polonium–210 (http:/ / hpschapters. org/ northcarolina/ NSDS/ 210PoPDF. pdf)". . Retrieved 2009-05-05. [35] " Effective half-life of polonium in the human (http:/ / www. osti. gov/ energycitations/ product. biblio. jsp?osti_id=7162390)". . Retrieved 2009-05-05. [36] " Polonium Poisoning (http:/ / nuclearweaponarchive. org/ News/ PoloniumPoison. html)". . Retrieved 2009-05-05. [37] Harrison J et al. (2007). "Polonium-210 as a poison". J. Radiol. Prot. 27: 17. doi: 10.1088/0952-4746/27/1/001 (http:/ / dx. doi. org/ 10. 1088/ 0952-4746/ 27/ 1/ 001). "The conclusion is reached that 0.1–0.3 GBq or more absorbed to blood of an adult male is likely to be fatal within 1 month. This corresponds to ingestion of 1–3 GBq or more, assuming 10% absorption to blood". [38] " National Academy of Sciences 1988 report: Health Risks of Radon and Other Internally Deposited Alpha-Emitters: BEIR IV, page 5 (http:/ / fermat. nap. edu/ openbook. php?record_id=1026& page=5)". .
10
Polonium Retrieved 2009-05-05. [39] " National Academy of Sciences 1999 report: Health Effects Of Exposure To Indoor Radon (http:/ / newton. nap. edu/ html/ beir6/ )". . Retrieved 2009-05-05. [40] " The Straight Dope: Does smoking organically grown tobacco lower the chance of lung cancer? (http:/ / www. straightdope. com/ columns/ 070928. html)". . Retrieved 2009-05-05. 210 [41] " Nuclear Regulatory Commission limits for Po (http:/ / www. nrc. gov/ reading-rm/ doc-collections/ cfr/ part020/ appb/ Polonium-210. html)". U.S. NRC. 2008-12-12. . Retrieved 2009-01-12. [42] " PilgrimWatch - Pilgrim Nuclear - Health Impact (http:/ / www. pilgrimwatch. org/ health1. html)". . Retrieved 2009-05-05. [43] Peter D. Zimmerman (2006). " The Smoky Bomb Threat (http:/ / www. nytimes. com/ 2006/ 12/ 19/ opinion/ 19zimmerman. html)". The New York Times. . Retrieved 2006-12-19. [44] " The mystery of Litvinenko's death (http:/ / news. bbc. co. uk/ 1/ hi/ uk/ 6180432. stm)". BBC News. 24 November 2006. . [45] " UK requests Lugovoi extradition (http:/ / news. bbc. co. uk/ 1/ hi/ uk/ 6698545. stm)". . Retrieved 2009-05-05. [46] " Focus: Cracking the code of the nuclear assassin (http:/ / www. timesonline. co. uk/ article/ 0,,2087-2484295_1,00. html)". . [47] " Restaurant Polonium: In Sheffield klingeln die Kassen (http:/ / www. zeit. de/ news/ artikel/ 2006/ 12/ 05/ 83406. xml)" (in German). Die Zeit (ZEIT online GmbH). 2006-12-05. . Retrieved 2008-06-06. [48] " Business booming at Polonium restaurant in English city, manager says (http:/ / www. iht. com/ articles/ ap/ 2006/ 12/ 01/ europe/ EU_GEN_Britain_Polonium_Restaurant. php)". International Herald Tribune. 2006-12-01. . Retrieved 2008-06-06. [49] " Why 'Polonium Restaurant' in UK is a hit (http:/ / timesofindia. indiatimes. com/ articleshow/ 712529. cms)". The Times of India. 2006-12-05. . [50] " Innocent chemical a killer (http:/ / www. news. com. au/ dailytelegraph/ story/ 0,22049,20863878-5001031,00. html)". The Daily Telegraph (Australia). December 4, 2006. . Retrieved 2009-05-05. [51] Karpin, Michael (2006). The bomb in the basement: How Israel went nuclear and what that means for the world. Simon and Schuster. ISBN 0743265947. [52] Maugh, Thomas; Karen Kaplan (2007-01-01). " A restless killer radiates intrigue (http:/ / articles. latimes. com/ 2007/ jan/ 01/ science/ sci-polonium1)". Los Angeles Times. . Retrieved 2008-09-17. [53] " Guidance for Industry. Internal Radioactive Contamination — Development of Decorporation Agents (http:/ / www. fda. gov/ downloads/ Drugs/ GuidanceComplianceRegulatoryInformation/ Guidances/ ucm071944. pdf)" (PDF). . Retrieved 2009-07-07. [54] Rencováa J., Svoboda V., Holuša R., Volf V., Jones M. M., Singh P. K. (1997). "Reduction of subacute lethal radiotoxicity of polonium-210 in rats by chelating agents". International Journal of Radiation Biology 72: 247. doi: 10.1080/095530097143338 (http:/ / dx. doi. org/ 10. 1080/ 095530097143338). [55] " The smoky bomb threat - Opinion - International Herald Tribune (http:/ / www. iht. com/ articles/ 2006/ 12/ 19/ opinion/ edzimmer. php)". . Retrieved 2009-05-05. [56] " Solutions to Static Problems (http:/ / www. amstat. com/ solutions/ staticmaster. html)". Amstat Industries. . Retrieved 2006-12-01. [57] " Static Eliminator (http:/ / www. osmolabstore. com/ OsmoLabPage. dll?BuildPage& 1& 1& 1005)". GE Osmonics' Labstore. . Retrieved 2006-12-01. [58] " Staticmaster Antistatic Products (http:/ / www. 2spi. com/ catalog/ photo/ statmaster. shtml)". SPI Supplies. . Retrieved 2007-08-29. [59] " General domestic licenses for byproduct material (http:/ / www. nrc. gov/ reading-rm/ doc-collections/ cfr/ part031/ full-text. html)". . Retrieved 2009-05-05. [60] Singleton, Don (November 28, 2006). " The Availability of polonium-210 (http:/ / donsingleton. blogspot. com/ 2006/ 11/ polonium-210. html)". . Retrieved 2006-11-29. [61] " UnitedNuclear Isotopes for sale over the Internet (http:/ / www. unitednuclear. com/ isotopes. htm)". . Retrieved 2007-03-19. [62] " Litvinenko affair: now the man who warned him poisoned too (http:/ / www. guardian. co. uk/ russia/ article/ 0,,1962354,00. html)". . Retrieved 2009-05-05. [63] Hooper, Rowan (13 December 2006). " Natural selections: Murder in the genes? Polonium, peacocks—and a dead spy (http:/ / search. japantimes. co. jp/ cgi-bin/ fe20061213rh. html)". The Japan Times Online. . Retrieved 2006-12-13. [64] Radford EP Jr, Hunt VR (1964). "Polonium 210: a volatile radioelement in cigarettes". Science 143: 247. doi: 10.1126/science.143.3603.247 (http:/ / dx. doi. org/ 10. 1126/ science. 143. 3603. 247).
11
Polonium [65] Kelley TF (1965). "Polonium 210 content of mainstream cigarette smoke". Science 149: 537. doi: 10.1126/science.149.3683.537 (http:/ / dx. doi. org/ 10. 1126/ science. 149. 3683. 537). [66] Hussein EM (1994). "Radioactivity of phosphate ore, superphosphate, and phosphogypsum in Abu-zaabal phosphate". Health Physics 67: 280. doi: 10.1097/00004032-199409000-00010 (http:/ / dx. doi. org/ 10. 1097/ 00004032-199409000-00010). [67] Barisic D, Lulic S, Miletic P (1992). "Radium and uranium in phosphate fertilizers and their impact on the radioactivity of waters". Water Research 26: 607. doi: 10.1016/0043-1354(92)90234-U (http:/ / dx. doi. org/ 10. 1016/ 0043-1354(92)90234-U).. [68] Scholten LC, Timmermans CWM (1992). "Natural radioactivity in phosphate fertilizers". Nutrient cycling in agroecosystems 43: 103. doi: 10.1007/BF00747688 (http:/ / dx. doi. org/ 10. 1007/ BF00747688). [69] Tidd J (2008). "The big idea: polonium, radon and cigarettes". Journal of the Royal Society of Medicine 101: 156. doi: 10.1258/jrsm.2007.070021 (http:/ / dx. doi. org/ 10. 1258/ jrsm. 2007. 070021). [70] William Birnbauer (September 7, 2008). " Big Tobacco covered up radiation danger (http:/ / www. theage. com. au/ national/ big-tobacco-covered-up-radiation-danger-20080906-4b54. html?page=-1)". The Age, Melbourne, Australia. . [71] http:/ / www. techlib. com/ science/ ion. html#PoloniumPen
12
Article Sources and Contributors
Article Sources and Contributors Polonium Source: http://en.wikipedia.org/w/index.php?oldid=308874415 Contributors: 007bond05, A bit iffy, AThing, Abune, Acroterion, AdamFunk, Aesopos, Ahoerstemeier, Alanevans, Alanjones25, Alansohn, Alexd17, Alexmcfire, Algocu, [email protected], AlimanRuna, Alphachimp, Alsandro, Anclation, Anna Britten, Arkuat, Astatine-210, Aste, Athipedia, Athomic69, AtomicFire, AussieBoy, Awesomebitch, AySz88, Baccyak4H, Badseed, Barnas, Baronnet, Baynardo7, Beetstra, Bender235, Benjah-bmm27, BerserkerBen, Bewildebeast, BillFlis, Biophys, BirdValiant, Bladeofgrass, BlueEarth, Bobo192, Bocianski, Borislav Dopudja, Brandon, Bryan Derksen, Bubbles34 11, Bucky242, By97aa, CP\M, CWY2190, CYD, Cadmium, Canadian-Bacon, Captain Video, Carnildo, Cgingold, Chameleon, Chaser, ChazYork, ChemNerd, ChirpingPenguin, Chivista, Chnv, Chris 73, Chuckwatson, Ciphergoth, Cjpuffin, Cmprince, Coemgenus, Conversion script, Coredesat, Corti, CrazyChemGuy, Crywalt, DB 8870, DGJM, DMacks, DalekThay, Damnreds, Dan100, Dancraggs, Dangfrick, Daniel Olsen, Darkl0rd182, Darth Panda, David Haslam, David Latapie, Davidruben, Deglr6328, DeltaF508, Dfrg.msc, Dina, Diomidis Spinellis, Dirac66, Discospinster, DocWatson42, Donreed, Dowcet, Downward machine, DragonflySixtyseven, Drbug, Dreish, Dricherby, Dulciana, Duncan.france, Dwchin, Dysprosia, Dzubint, ESkog, Ed Fitzgerald, Ed4linda, Edgar181, Editor at Large, El C, Eleuther, Elmarco, Emc2, Emmiep, Emperorbma, Epbr123, Eridani, Erik Zachte, Erth64net, Esoltas, Essent, Faiellie, Feline1, Femto, Ferkelparade, Fibonacci, Floaterfluss, Fodio, Freticat, Fungible, Fvasconcellos, Fys, GDonato, Gabi S., Gaius Cornelius, Gareth.randall, Garion96, Gdavidp, Gene Nygaard, Giftlite, Gloop, Gogetinaditchyouhobo, Goldfinch, Goudzovski, GracieLizzie, Grahamers, Gritzko, Gugilymugily, Gurch, HMGriffin, HPaul, Hankwang, HappyCamper, HazyM, Headybrew, Hektor, Herkdrvr, Hqb, Humanist, IForgotToEatBreakFast, IW.HG, Icairns, Ideyal, Imjustmatthew, Itai, J.delanoy, Jackbonneruk, Janke, Jao, Jaraalbe, Jauerback, Jeepday, Jeepien, Jeff G., Jellyfisho, Jeltz, Jimp, Jinma, Jkl, Jmount, Joanjoc, Joegee, John, Joseph 2166, Juansmith, Julesd, Jurand, Justicio, Karn, Kay Dekker, Kenosis, Kerowren, Kev-o1102, Kghusker, Khalid hassani, Khukri, Kirils, Kman90, Kpalion, Kpjas, KudzuVine, Kwamikagami, Kwertii, Lawpuh, LeaveSleaves, Lennylim, LibLord, LilHelpa, Lincmad, Lklundin, Logicgate00, LokiiT, Lolitaona, Lord Anubis, LorenzoB, Lost tourist, Lradrama, LukeSurl, MWS, Madder, Man with two legs, Marc Venot, Marielle24, MarkSutton, Markco1, Martin451, Materialscientist, Mathiasrex, Mausy5043, Mav, McSly, Mdf, Meertn, Mglg, Mieciu K, MightyWarrior, Minesweeper, Mistercow, Misza13, Mondebleu, Morwen, Moshe Constantine Hassan Al-Silverburg, Muttonking, Mycroft.Holmes, Mysterychopper3027, NSH001, Nabokov, Nate Biggs, NawlinWiki, NcSchu, Nergaal, Nick, Nick Y., NickBush24, Niffux, Nivagh, Nrcprm2026, Nyq, Nyttend, Opirt, Origamikid, Oxymoron83, PBIPhotobug, PGWG, Paraballo, Pasachoff, Patrick, PeterJeremy, Petri Krohn, Pgk, Philip Trueman, Piano non troppo, Pilotguy, Pkoppenb, Plasticup, PlatinumX, Pleonic, Pol098, Polonium, Poolkris, Popo le Chien, Pqrstuv, Rcnet, Remember, Rifleman 82, Right angle, Rjwilmsi, Rolypolyman, Romanm, Rory096, Royalbroil, Rrburke, Rubicon, Rundquist, Ryan Reich, Rydra Wong, SWAdair, Sanchom, Saperaud, Schneelocke, Sengkang, Shanes, Shinkolobwe, Shultzc, Sisu99, Sl, Sladen, Slogby, Sofakingon, Sotakeit, Spiel496, Srleffler, Srock-rocks, StealthFox, Steel, StephP, Stephen B Streater, Stephenb, Stifynsemons, Stone, Stui, Svante, SweetNeo85, Swid, Symon, SynergyBlades, T-borg, Talon Artaine, Tapir Terrific, Tdoublenineone, Tetracube, The Cunctator, The Font, The Rambling Man, Thefool12367, Theo Pardilla, Thingg, TimVickers, Timothylord, Tinsman, Tjmayerinsf, Trumpet marietta 45750, Tsp110, Until It Sleeps, V1adis1av, VTNC, Versus22, Vgranucci, Vsmith, Wanted, Watch37264, Welsh, What!?Why?Who?, WhiteCurse15, Wikifier, Wikikids07, Willie xD, Witkacy, Wtmitchell, Wtwjgc, Xaje, Xy7, YBeayf, Yaf, Yekrats, Yerpo, Yyy, ZayZayEM, Zfr, Zoporific, Zowie, Zzuuzz, 561 anonymous edits
Image Sources, Licenses and Contributors image:Po-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Po-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Paddy, Saperaud Image:alpha po lattice.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Alpha_po_lattice.jpg License: Public Domain Contributors: Original uploader was Cadmium at en.wikipedia Image:Gammaspectrscopy.png Source: http://en.wikipedia.org/w/index.php?title=File:Gammaspectrscopy.png License: Public Domain Contributors: Bryan Derksen, Cadmium, MithrandirMage, 3 anonymous edits Image:Alpha1spec.png Source: http://en.wikipedia.org/w/index.php?title=File:Alpha1spec.png License: Public Domain Contributors: Cadmium, MithrandirMage, 1 anonymous edits Image:Alpha5spec.png Source: http://en.wikipedia.org/w/index.php?title=File:Alpha5spec.png License: Public Domain Contributors: Cadmium, MithrandirMage, 1 anonymous edits
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
13
Astatine
1
Astatine polonium ← astatine → radonI ↑ At ↓ Uus
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
85At Periodic table
Appearance black solid (presumed) General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseMelting pointBoiling pointHeat of vaporizationVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
361
392
429
475
531
607
Atomic properties Oxidation states ElectronegativityIonization energiesCovalent radius Miscellaneous Magnetic orderingThermal conductivityCAS registry number Most stable isotopes Main article: Isotopes of astatine iso 210
At
211
At
N.A. syn
syn
half-life 8.1 h
DM
DE (MeV)
DP
ε, β+
3.981
210
α
5.631
206
Po Bi
7.2 h
astatine, At, 85 halogens17, 6, p(210) g·mol−1 [Xe] 4f14 5d10 6s2 6p5 2, 8, 18, 32, 18, 7 (Image) solid 575 K,302 °C,576 °F 610 K,337 °C,639 °F 40 kJ·mol−1 ±1, 3, 5, 7 2.2 (Pauling scale) 1st: 890±40 kJ·mol−1 150 pm no data (300 K) 1.7 W·m−1·K−1 7440-68-8
Astatine Astatine (pronounced /ˈæstətiːn/ or /ˈæstətɨn/) is a radioactive chemical element with the symbol At and atomic number 85. It is the heaviest of the discovered halogens. Although astatine is produced by radioactive decay in nature, due to its short half life it is found only in minute amounts. Astatine was first produced by Dale R. Corson, Kenneth Ross MacKenzie, and Emilio Segrè in 1940. Three years passed before traces of astatine were also found in natural minerals. Until recently most of the physical and chemical characteristics of astatine were inferred from comparison with other elements. Some astatine isotopes are used as alpha-particle emitters in science applications, and medical applications for astatine 211 have been tested. Astatine is currently the rarest naturally-occurring element, with less than an ounce contained in the entire Earth's crust.[1]
Characteristics This highly radioactive element has been confirmed by mass spectrometers to behave chemically much like other halogens, especially iodine (it would probably accumulate in the thyroid gland like iodine), though astatine is thought to be more metallic than iodine. Researchers at the Brookhaven National Laboratory have performed experiments that have identified and measured elementary reactions that involve astatine;[2] however, chemical research into astatine is limited by its extreme rarity, which is a consequence of its extremely short half-life. Its most stable isotope has a half-life of around 8.3 hours. The final products of the decay of astatine are isotopes of lead. The halogens get darker in color with increasing molecular weight and atomic number. Thus, following the trend, astatine would be expected to be a nearly black solid, which, when heated, sublimes into a dark, purplish vapor (darker than iodine). Astatine is expected to form ionic bonds with metals such as sodium, like the other halogens, but it can be displaced from the salts by lighter, more reactive halogens. Astatine can also react with hydrogen to form hydrogen astatide, which when dissolved in water, forms the exceptionally strong hydroastatic acid. Astatine is the least reactive of the halogens, being less reactive than iodine.[3]
History The existence of "eka-iodine" had been predicted by Mendeleev. Astatine (after Greek αστατος astatos, meaning "unstable") was first synthesized in 1940 by Dale R. Corson, Kenneth Ross MacKenzie, and Emilio Segrè at the University of California, Berkeley by bombarding bismuth with alpha particles.[4] As the periodic table of elements was long known, several scientists tried to find the element following iodine in the halogen group. The unknown substance was called Eka-iodine before its discovery because the name of the element was to be suggested by the discoverer. The claimed discovery in 1931 at the Alabama Polytechnic Institute (now Auburn University) by Fred Allison and associates, led to the spurious name for the element as alabamine (Ab) for a few years.[5] [6] [7] This discovery was later shown to be an erroneous one. The name Dakin was proposed for this element in 1937 by the chemist Rajendralal De working in Dhaka, Bangladesh.[8] The name Helvetium was chosen by the Swiss chemist Walter Minder, when he announced the discovery of element 85 in 1940, but changed his suggested name to Anglohelvetium in 1942.[9]
2
Astatine
3
It took three years before astatine was found as product of the natural decay processes. The short-lived element was found by the two scientists Berta Karlik and Traude Bernert.[10] [11]
Occurrence Astatine occurs naturally in three natural radioactive decay series, but because of its short half-life is found only in minute amounts. Astatine-218 (218At) is found in the uranium series, 216At is in the thorium series, and 215At as well as 219At are in the actinium series[12] . The most long-lived of these naturally-occurring astatine isotopes is 219At with a half-life of 56 seconds. Astatine is the rarest naturally-occurring element, with the total amount in Earth's crust estimated to be less than 1 oz (28 g) at any given time. This amounts to less than one teaspoon of the element. Guinness World Records has dubbed the element the rarest on Earth, stating: "Only around 0.9 oz (25 g) of the element astatine (At) occurring naturally". Isaac Asimov, in a 1957 essay on large numbers, scientific notation, and the size of the atom, wrote that in "all of North and South America to a depth of ten miles", the number of astatine atoms at any time is "only a trillion".[13]
Production for alpha particles of 26 MeV[14] for alpha particles of 40 MeV[14] for alpha particles of 60 MeV.[15] Astatine is produced by bombarding bismuth with energetic alpha particles to obtain relatively long-lived 209At - 211At, which can then be distilled from the target by heating in the presence of air.
Compounds Multiple compounds of astatine have been synthesized in microscopic amounts and studied as intensively as possible before their inevitable radioactive disintegration. The reactions are normally tested with dilute solutions of astatine mixed with larger amounts of iodine. The iodine acts as a carrier, ensuring that there is sufficient material for laboratory techniques such as filtration and precipitation to work.[14] [16] While these compounds are primarily of theoretical interest, they are being studied for potential use in nuclear medicine.[17] Astatine is expected to form ionic bonds with metals such as sodium, like the other halogens, but it can be displaced from the salts by lighter, more reactive halogens. Astatine can also react with hydrogen to form hydrogen astatide (HAt), which when dissolved in water, forms hydroastatic acid. Some examples of astatic compounds are: NaAt or sodium astatide MgAt2 or magnesium astatide CAt4 or carbon tetrastatide (tetraastatide)
Astatine
4
Isotopes Astatine has 33 known isotopes, all of which are radioactive; the range of their mass numbers is from 191 to 223. There exist also 23 metastable excited states. The longest-lived isotope is 210At, which has a half-life of 8.1 hours; the shortest-lived known isotope is 213At, which has a half-life of 125 nanoseconds.[18]
Applications The least stable isotopes of astatine have no practical applications other than scientific study due to their extremely short life, but heavier isotopes have medical uses. Astatine 211 is an alpha emitter with a physical half-life of 7.2 h. These features have led to its use in radiation therapy.[19] An investigation of the efficacy of astatine-211—tellurium colloid for the treatment of experimental malignant ascites in mice reveals that this alpha-emitting radiocolloid can be curative without causing undue toxicity to normal tissue. By comparison, beta-emitting phosphorus-32 as colloidal chromic phosphate had no antineoplastic activity. The most compelling explanation for this striking difference is the dense ionization and short range of action associated with alpha-emission. These results have important implications for the development and use of alpha-emitters as radiocolloid therapy for the treatment of human tumors.[20]
Precautions Since astatine is extremely radioactive, it should be handled with extreme care. Because of its extreme rarity, it is not likely that the general public will be exposed. Astatine is a halogen, and standard precautions apply. It is reactive, sharing similar chemical characteristics with iodine. There are toxicologic studies of astatine-211 on mice indicating that radioactive poisoning is the major effect on living organisms. [21]
External links • WebElements.com - Astatine [22] • Doc Brown's Chemistry Clinic - Group 7 The Halogens
[23]
pnb:نیٹاٹسیا
References [1] Close, Frank. Particle Physics: A Very Short Introduction. Oxford University Press: New York, 2004. Page 2. [2] C. R. Hammond. The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304857. [3] Anders, E. (1959). "Technetium and Astatine Chemistry". Annual Review of Nuclear Science 9: 203–220. doi: 10.1146/annurev.ns.09.120159.001223 (http:/ / dx. doi. org/ 10. 1146/ annurev. ns. 09. 120159. 001223). [4] D. R. Corson, K. R. MacKenzie, and E. Segrè (1940). "Artificially Radioactive Element 85". Phys. Rev. 58: 672–678. doi: 10.1103/PhysRev.58.672 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 58. 672). [5] Fred Allison, Edgar J. Murphy, Edna R. Bishop, and Anna L. Sommer (1931). "Evidence of the Detection of Element 85 in Certain Substances". Phys. Rev. 37: 1178–1180. doi: 10.1103/PhysRev.37.1178 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 37. 1178). [6] " Alabamine & Virginium (http:/ / www. time. com/ time/ magazine/ article/ 0,9171,743159,00. html)". time. . Retrieved 2008-07-10. [7] Trimble, R. F. (1975). "What happened to alabamine, virginium, and illinium?". J. Chem. Educ. 52: 585. [8] 85 Astatine (http:/ / elements. vanderkrogt. net/ elem/ at. html)
Astatine [9] Alice Leigh-Smith, Walter Minder (1942). "Experimental Evidence of the Existence of Element 85 in the Thorium Family". Nature 150: 767–768. doi: 10.1038/150767a0 (http:/ / dx. doi. org/ 10. 1038/ 150767a0). [10] Karlik, Berta; Bernert Traude (1943). "Eine neue natürliche α-Strahlung". Naturwissenschaften 31 (25–26): 298–299. doi: 10.1007/BF01475613 (http:/ / dx. doi. org/ 10. 1007/ BF01475613). [11] Karlik, Berta; Bernert Traude (1943). "Das Element 85 in den natürlichen Zerfallsreihen". Zeitschrift für Physik 123 (1–2): 51–72. doi: 10.1007/BF01375144 (http:/ / dx. doi. org/ 10. 1007/ BF01375144). [12] " astatine (At) (http:/ / www. britannica. com/ eb/ article-9009963/ astatine)". Encyclopedia Britannica online. . Retrieved 2008-06-22. [13] http:/ / ia331335. us. archive. org/ 1/ items/ onlyatrillion017765mbp/ onlyatrillion017765mbp_djvu. txt [14] Nefedov, V. D. (1968). "Astatine". Russ. Chem. Rev. 37: 87–98. doi: 10.1070/RC1968v037n02ABEH001603 (http:/ / dx. doi. org/ 10. 1070/ RC1968v037n02ABEH001603). [15] Barton, G. W.; Ghiorso, A.; Perlman, I. (1951). "Radioactivity of Astatine Isotopes". Physical Reviews 82 (1): 13–19. doi: 10.1103/PhysRev.82.13 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 82. 13). [16] Aten Jun., A. H. W.; Doorgeest, T.; Hollstein U.; Moeken H. P. (1952). "Section 5: radiochemical methods. Analytical chemistry of astatine". Analyst 77: 774–777. doi: 10.1039/AN9527700774 (http:/ / dx. doi. org/ 10. 1039/ AN9527700774). [17] Boyd, Jade (1007-08-27). " Nuclear Nanocapsules, The New Cancer Weapon (http:/ / www. medicalnewstoday. com/ articles/ 80581. php)". Medical News Today. . Retrieved 2008-11-05. [18] Audi, Georges (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3–128. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [19] Wilbur, D. S. (01 Oct 2001). " Overcoming the Obstacles to Clinical Evaluation of 211At-Labeled Radiopharmaceuticals. (http:/ / jnm. snmjournals. org/ cgi/ reprint/ 42/ 10/ 1516)". Journal Nuclear Medicine 42 (10): 1516–1518. . [20] Bloomer, W. D.; McLaughlin, W. H.; Neirinckx, R. D.; Adelstein, S. J.; Gordon, P. R.; Ruth, T. J.; Wolf, AP (1981). " Astatine-211--tellurium radiocolloid cures experimental malignant ascites (http:/ / www. sciencemag. org/ cgi/ content/ abstract/ 212/ 4492/ 340)". Science 212 (4492): 340–341. doi: 10.1126/science.7209534 (http:/ / dx. doi. org/ 10. 1126/ science. 7209534). PMID 7209534. . [21] Cobb, L.M. (1988). "Toxicity of Astatine-211 in the Mouse". Human & Experimental Toxicology 7: 529. doi: 10.1177/096032718800700602 (http:/ / dx. doi. org/ 10. 1177/ 096032718800700602). [22] http:/ / www. webelements. com/ webelements/ elements/ text/ At/ index. html [23] http:/ / www. docbrown. info/ page03/ The_Halogens. htm
5
Article Sources and Contributors
Article Sources and Contributors Astatine Source: http://en.wikipedia.org/w/index.php?oldid=308468438 Contributors: !Astatine210isalo$er, Ahoerstemeier, AlimanRuna, Altenmann, AmosJLedge, Anclation, Andre Engels, Andres, Andrew Kanode, Arakunem, Arkuat, Aurocker49, Awesome Truck Ramp, Axl, Beetstra, Benbest, Bentley4, BlueEarth, Borameer, Brianski, Bryan Derksen, Bsadowski1, Bulbear 287, Bullzeye, CLoWnnSkuLL, CYD, Caltas, CamperStrike, Can't sleep, clown will eat me, Caniemiec, Canley, CapitalR, Capricorn42, Cesium 133, Colbuckshot, Conversion script, Cryptic C62, Cyde, DMacks, Dadude3320, Dajwilkinson, Danj729, David Latapie, Delldot, Dendodge, DocWatson42, Donarreiskoffer, DrBob, DragonflySixtyseven, Dysepsion, Edgar181, El C, Emperorbma, Erik Zachte, Erockrph, Esrever, Euchiasmus, Farosdaughter, Femto, FocalPoint, Frencheigh, Gamingexpert, Gogo Dodo, Haikupoet, Hammer1980, Hashar, Heron, Hqb, Husond, Icairns, Ideyal, Itub, JRM, Jaraalbe, Jeff G., Jiang, Joe Masterguns, Kalamkaar, Kay Dekker, Kbdank71, Keilana, Kingdon, Kiwi137, Koyaanis Qatsi, Kupos, Kurykh, Kwamikagami, Landsword01, LarryMorseDCOhio, Legionnaire1718, Lightmouse, Mac Davis, Machete97, Marc Venot, Mashford, Materialscientist, Mav, Mdebets, Minesweeper, Mr. Lefty, Mr0t1633, Natural Cut, Nergaal, NewEnglandYankee, Nihiltres, No1lakersfan, OlEnglish, Omexis, Oxymoron83, PP Jewel, Paddu, PeterJeremy, Phrodu, Physchim62, Piperh, PlatinumX, Pol098, Polonium, Poolkris, Pras, Quadell, Regardless143, Remember, Rettetast, Reyk, Reza kalani, Rjwilmsi, Roberta F., Roentgenium111, Romanm, Rune.welsh, Rxnd, SLi, Sam Hocevar, Saperaud, Sbharris, Scarian, Schneelocke, SeanSchricker, Sengkang, Sharkface217, SidP, Silvonen, Sl, SnetskyCM, Snowmanmelting, Speciman00, Sponisdude, Stone, The Evil Spartan, The-chemist, Thingg, Titoxd, Toplegochamp, Trevor.tombe, V1adis1av, Van helsing, Vsmith, Vuerqex, Vuong Ngan Ha, Wandering Courier, Wii Wiki, XIDE, Xenophon777, Xy7, Yamamoto Ichiro, Yekrats, Yuanchosaan, Yyy, Zachwoo, Zigger, 263 anonymous edits
Image Sources, Licenses and Contributors file:Unknown.svg Source: http://en.wikipedia.org/w/index.php?title=File:Unknown.svg License: Public Domain Contributors: Mav file:Electron shell 085 Astatine.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_085_Astatine.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
6
Radon
1
Radon astatine ← radon → franciumXe ↑ Rn ↓ Uuo
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
86Rn Periodic table
Appearance colorless General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseMelting pointBoiling pointCritical pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
at T/K
110
121
134
152
176
211
Atomic properties Oxidation states ElectronegativityIonization energiesCovalent radius Miscellaneous Crystal structureMagnetic orderingThermal conductivityCAS registry number Most stable isotopes Main article: Isotopes of radon iso
N.A.
half-life
DM
DE (MeV)
DP
210
syn
2.4 h
Alpha
6.404
216
211
syn
14.6 h
Epsilon
2.892
211
Alpha
5.965
207
Rn Rn
Po At Po
222
trace
3.8235 d
Alpha
5.590
218
224
syn
1.8 h
Beta
0.8
224
Rn Rn
Po Fr
Radon
radon, Rn, 86 noble gases 18, 6, p(222) g·mol−1 [Xe] 4f14 5d10 6s2 6p6 2, 8, 18, 32, 18, 8 (Image) gas 202 K,−71.15 °C,−96 °F 211.3 K,−61.85 °C,−79.1 °F 377 K, 6.28 MPa 3.247 kJ·mol−1 18.10 kJ·mol−1 (25 °C) 20.786 J·mol−1·K−1 2 2.2 (Pauling scale) 1st: 1037 kJ·mol−1 150 pm face-centered cubic non-magnetic (300 K) 3.61 m W·m−1·K−1 10043-92-2 Radon (pronounced /ˈreɪdɒn/) is a chemical element with symbol Rn and atomic number 86. Radon is a colorless, odorless, tasteless, naturally occurring, radioactive noble gas that is formed from the decay of radium. It is one of the heaviest substances that remains a gas under normal conditions and is considered to be a health hazard. The most stable isotope, 222 Rn, has a half-life of 3.8 days. While having been less studied by chemists due to its high radioactivity, there are a few known compounds of this generally unreactive element. Radon is formed from the normal radioactive decay of uranium. Uranium has been around since the earth was formed and its most common isotope has a very long half-life (4.5 billion years), which is the amount of time required for one-half of uranium to break down. Uranium, radium, and thus radon, will continue to exist indefinitely at about the same levels as they do now.[1] Radon is responsible for the majority of the mean public exposure to ionizing radiations. It is often the single largest contributor to an individual's background radiation dose, and is certainly the most variable from location to location. Radon gas from natural sources can accumulate in buildings, especially in confined areas such as basements. Radon can be found in some spring waters and hot springs.[2] Breathing high concentrations of radon can cause lung cancer. Thus, radon is considered a significant contaminant that affects indoor air quality worldwide. According to the United States Environmental Protection Agency, radon could be the second most frequent cause of lung cancer, after cigarette smoking; and radon-induced lung cancer the 6th leading cause of cancer death overall, causing 21,000 lung cancer deaths per year in the United States.[3]
2
Radon
3
History and etymology Discovered in 1900 by Friedrich Ernst Dorn, radon was the fifth radioactive element to be discovered, after uranium, thorium, radium and polonium.[4] [5] [6] In 1900 Dorn reported some experiments in which he noticed that radium compounds emanate a radioactive gas which he named Radium Emanation (Ra Em).[7] Before that, in 1899, Pierre and Marie Curie observed that the "gas" emitted by radium remained radioactive for a month.[8] Later that year, Robert B. Owens and Ernest Rutherford noticed variations when trying to measure radiation from thorium oxide.[9] Rutherford noticed that the compounds of thorium continuously emit a radioactive gas which retain the radioactive powers for several minutes and called this gas "emanation" (from Latin "emanare"—to elapse and "emanatio"—expiration),[10] and later Thorium Emanation (Th Em). In 1901, he demonstrated that the emanations are radioactive, but credited the Curies for the discovery of the element.[11] In 1903, similar emanations were observed from actinium by André-Louis Debierne[12] [13] and were called Actinium Emanation (Ac Em). Several names were suggested for these three gases: exradio, Apparatus used by Ramsay and Whytlaw-Gray to isolate radon. M is a capillary tube where approximately 0.1 mm3 were isolated. Rn mixed with H2 entered the evacuated system through siphon A; mercury is shown in black.
exthorio, and exactinio in 1904;[14] radon, thoron, and akton in 1918;[15] radeon, thoreon, and actineon in 1919,[16] and eventually radon, thoron, and actinon in 1920.[17] The likeness of the spectra of these three gases with those of argon, krypton, and xenon, and their observed chemical inertia led Sir William Ramsay to suggest in 1904 that the "emanations" might contain a new element of the noble gas family.[14]
In 1910, Sir William Ramsay and Robert Whytlaw-Gray isolated radon, determined its density, and determined that it was the heaviest known gas.[18] They wrote that "L'expression de l'émanation du radium est fort incommode," (the expression of radium emanation is very awkward) and suggested the new name niton (Nt) (from the Latin "nitens" meaning "shining") in order to emphasize the property of gases that cause the phosphorescence of some substances,[18] and in 1912 it was accepted by the International Commission for Atomic Weights. In 1923, the International Committee for Chemical Elements and International Union of Pure and Applied Chemistry (IUPAC) chose among the names radon (Rn), thoron (Tn), and actinon (An). Later, when isotopes were numbered instead of named, the element took the name of the most stable isotope, radon, while Tn became 220Rn and An 219Rn. As late as the 1960s, the element was also referred to simply as emanation.[19] The first synthesized compound of radon, radon fluoride, was obtained in 1962.[20] The danger of high exposure to radon in mines, where exposures reaching 1,000,000 Bq/m3 can be found, has long been known. In 1530, Paracelsus described a wasting disease of miners, the mala metallorum, and Georg Agricola recommended ventilation in mines to avoid this mountain sickness (Bergsucht).[21] [22] In 1879, this condition was identified as
Radon lung cancer by Herting and Hesse in their investigation of miners from Schneeberg, Germany. The first major studies with radon and health occurred in the context of uranium mining, first in the Joachimsthal region of Bohemia and then in the Southwestern United States during the early Cold War.
Characteristics Physical form Radon is a colorless and odorless gas, and therefore not readily detectable by human senses alone. At standard temperature and pressure, radon forms a monatomic gas with a density of 9.73 kg/m3,[23] about 8 times the surface density of the Earth's atmosphere, 1.217 kg/m3,[24] and is one of the heaviest gases at room temperature and the heaviest of the noble gases, excluding ununoctium. At standard temperature and pressure, radon is a colorless gas, but when it is cooled below its freezing point of 202 K (−71 °C; −96 °F), it has a brilliant phosphorescence which turns yellow as the temperature is lowered, and becomes orange-red as the air liquefies at temperatures below 93 K (−180.1 °C; −292.3 °F).[25] Upon condensation, radon also glows because of the intense radiation it produces.[26] Because it is radioactive and is a relatively unreactive chemical element, radon has few uses and is seldom used in academic research.
Isotopes Radon has no stable isotopes. There are 36 radioactive isotopes that have been characterized which range from an atomic mass of 193 to 228.[27] The most stable isotope is 222Rn, which is a decay product of 226Ra. It has a half-life of 3.823 days and decomposes by alpha particle emission into 218Po.[27] Among the decay daughters of this decay chain is also the highly unstable isotope 218Rn. The naturally occurring 226Ra is a product of the decay chain of 238U.[28] (The Wikipedia article on decay chain of 238U lists all the decay products of 222Rn.) There are three other isotopes that have a half life of over an hour: 211Rn, 210Rn and 224Rn. The 220Rn isotope is a natural decay product of the most stable thorium isotope (232Th), named thoron. It has a half-life of 55.6 seconds and also emits alpha radiation. Similarly, 219 Rn is derived from the most stable isotope of actinium (227Ac)—named “actinon”—and is an alpha emitter with a half-life of 3.96 seconds.[27] No radon isotopes are part of the other major decay series, that of neptunium (237Np).
4
Radon
5
Chemistry Radon is a member of the zero-valence elements that are called noble or inert gases. It is inert to most common chemical reactions, such as combustion, because the outer valence shell contains eight electrons. This produces a stable, minimum energy configuration in which the outer electrons are tightly bound.[29] For example, energy of more than 248 kcal/mol is required to extract one electron from its shells (also known as the first ionization energy).[30] However, due to periodic trends, radon has a lower electronegativity than the element one period before it, xenon, and is therefore more reactive. Radon is sparingly soluble in water, but more soluble than An electron shell diagram for radon. Note the eight electrons in the outer lighter noble gases. Radon is appreciably more soluble shell. in organic liquids than in water. Early studies concluded that the stability of radon hydrate should be of the same order as that of the hydrates of chlorine (Cl2) or sulfur dioxide (SO2), and significantly higher than the stability of the hydrate of hydrogen sulfide (H2S).[31] Because of its price and radioactivity, experimental chemical research is seldom performed with radon, and as a result there are very few reported compounds of radon, all either fluorides or oxides. Radon can be oxidized by a few powerful oxidizing agents such as F2, thus forming radon fluoride.[32] [33] It decomposes back to elements at a temperature of above 250 °C. It has a low volatility and was thought to be RnF2. But because of the short half-life of radon and the radioactivity of its compounds, it has not been possible to study the compound in any detail. However, theoretical studies on this molecule have predicted that it should have a Rn-F bond distance of 2.08 Ǻ, and that the compound is thermodynamically more stable and less volatile than its lighter counterpart XeF2.[34] The octahedral molecule RnF6 was predicted to have an even lower enthalpy of formation than the difluoride.[35] The [RnF]+ is believed to form by the reaction:[36] Rn (g) + 2 [O2]+[SbF6]− (s) → [RnF]+[Sb2F11]− (s) + 2 O2 (g) Radon oxides are among the few other reported compounds of radon.[37] The radon carbonyl RnCO has been predicted to be stable and to have a linear geometry.[38] The molecules Rn2 and RnXe were found to be significantly stabilized by spin-orbit coupling.[39] Radon caged inside a fullerene has been proposed as a drug for tumors.[40]
Radon
6
Equilibrium factor with radon progenies 222
Rn belongs to the Radium and Uranium-238 decay chain. It decays with a half-life of 3.8235 days. Its four first progenies (excluding marginal decay schemes) are very short-lived, meaning that the corresponding disintegrations are correlated to the initial radon distribution : • • • •
218
Po, 3.10 minutes, alpha decay. Pb, 26.8 min, beta decay. 214 Bi, 19.9 min, beta decay. 214 Po, 0.1643 ms, alpha decay. 214
At the next step,
214
Po decays to
210
Pb, which has a much longer half-life of 22.3 years.
[41]
The radon equilibrium factor is the ratio between the activity of all short-period radon progenies (which are responsible of most of the biological effect), and the activity that would be at equilibrium with the radon parent. The equilibrium factor is 1 when both activities are equal, meaning that the decay products have stayed close to the radon parent, long enough for the equilibrium to be reached (a couple of hours). If a closed volume is constantly supplied with radon, the concentration of short-lived daughters will increase until an equilibrium is reached where the rate of decay of each daughter will equal that of the radon itself. Under these conditions each pCi/l of radon will give rise to (almost precisely) 0.01 WL (see explanation of WL below). Ordinarily these conditions do not hold: in homes, the equilibrium fraction is typically 40%; that is, there will be 0.004 WL of progeny for each pCi/l of radon in air.[42] Progenies adhere to objects or dust particles (because of their electrostatic charge), whereas gaseous radon does not, so that the equilibrium factor in the atmosphere is usually less than one. The equilibrium factor is lowered by air circulation or air filtration devices. It is increased by air borne dust particles (such as cigarette smoke). The equilibrium factor retained in epidemiological studies is 0.4.[43] 210
Pb has a half-life of 22.3 years. Its progenies are:
• • •
210
Bi, 5.013 days, beta decay. Po, 138.376 days, alpha decay. 206 Pb, stable. 210
Radon
Occurrence Concentration units All discussions of radon concentrations in the environment refer to 222Rn. While the average rate of production of 220Rn (from the thorium decay series) is about the same as 222Rn, the amount of 220 Rn entering the environment is much less than that of 222Rn because of the short half-life of 220 Rn (1 minute versus 4 days).[1] Radon concentrations found in 210 natural environments are much Pb is formed from the decay of 222Rn. Here is a typical deposition rate of 210Pb as observed in Japan as a function of too low to be detected by chemical [44] time, due to variations in radon concentration. means. A 1000 Bq/m3 (relatively high) concentration corresponds to 0.17 pico-gram per cubic meter. The average concentration of radon in the atmosphere is about 6 × 10-20 atoms of radon for each molecule in the air, or about 150 atoms in each ml of air.[45] All the radon activity of the Earth atmosphere is due to some tens of grams of radon.[46] Radon concentration is usually measured in the atmosphere, in becquerel per cubic meter (Bq/m3), the SI derived unit. Typical domestic exposures are of ≈ 100 Bq/m3 indoors, and 10-20 Bq/m3 outdoors. It is often measured in pico-curie per liter (pCi/l) in the USA, with 1 pCi/l=37 Bq/m3 (or 5 pCi/l ≈ 200 Bq/m3).[42] In the mining industry, the exposition is traditionally measured in working level (WL), and the cumulative exposition in working level month (WLM) : 1 WL equals any combination of short-lived 222Rn progeny (218Po, 214Pb, 214Bi, and 214Po) in 1 liter of air that releases 1.3 x 105 MeV of potential alpha energy;[42] one WL is equivalent to 2.08 x 10^5 joules per cubic meter of air (J/m3).[1] The SI unit of cumulative exposure is expressed in joule-hours per cubic meter (J·h/m3). One WLM is equivalent to 3.6 x 10−3 J·h/m3. An exposure to 1 WL for 1 working month (170 hours) equals 1 WLM cumulative exposure. A cumulative exposition of 1 WLM is roughly equivalent to living one year in an atmosphere with a radon concentration of 230 Bq/m3.[47] The radon (222Rn) released into the air decays to 210Pb and other radioisotopes, the levels of 210Pb can be measured. The rate of deposition of this radioisotope is dependent on the weather.
7
Radon
Natural Radon is a decay product of uranium, which is relatively common in the Earth's crust, but generally concentrated in ore-bearing rocks scattered around the world. Every square mile of surface soil, to a depth of 6 inches (2.6 km2 to a depth of 15 cm), contains approximately 1 gram of radium, which releases radon in small amounts to the [1] atmosphere On a global scale, it is estimated that 2,400 million curies of radon are released from soil annually. Radon concentration varies wildly from place to place. In the open air, it ranges from 1 to 100 Bq/m3, even less (0.1 Bq/m3) above the ocean. In Radon concentration next to a uranium caves or aerated mines, or ill-aerated houses, its mine concentration climbs to 20-2,000 Bq/m3.[48] Radon concentration can be much higher in mining contexts. Due to ventilation regulation, typical radon concentration in uranium mines is usually maintained under the "working level", with 95th percentile levels ranging up to nearly 3 WL (546 pCi 222Rn per liter of air; 20,202 Bq/m3, measured from 1976 to 1985).[1] The concentration in the air at the (unventilated) Gastein Healing Gallery averages 43 kBq/m3 (about 1.2 nCi/l) with maximal value of 160 kBq/m3 (about 4.3 nCi/l).[49] Radon mostly appears with the decay chain of the radium and uranium series (222Rn), and marginally with the thorium series (220Rn). The element emanates naturally from the ground all over the world, wherever traces of uranium or thorium can be found, and particularly in regions with soils containing granite or shale, which have a higher concentration of uranium. However, not all granitic regions are prone to high emissions of radon. Being a rare gas, it usually migrates freely through faults and fragmented soils, and 222 may accumulate in caves or water. Due to its very small half-life (four days for Rn), its concentration decreases very quickly when the distance from the production area increases. Its concentration varies greatly with season and atmospheric conditions. For instance, it has been shown to accumulate in the air if there is a meteorological inversion and little wind.[50] High concentrations of radon can be found in some spring waters and hot springs.[51] The towns of Boulder, Montana; Misasa; Bad Kreuznach, Germany; and the country of Japan have radium-rich springs which emit radon. To be classified as a radon mineral water, radon concentration must be above a minimum of 2 nCi/l (74 Bq/l).[52] The activity of radon mineral water reaches 2,000 Bq/l in Merano and 4,000 Bq/l in Lurisia (Italy).[49] Natural radon concentrations in Earth's atmosphere are so low that radon-rich water in contact with the atmosphere will continually lose radon by volatilization. Hence, ground water has a higher concentration of 222Rn than surface water, because the radon is continuously produced by radioactive decay of 226Ra present in rocks. Likewise, the saturated zone of a soil frequently has a higher radon content than the unsaturated zone because of diffusional losses to the atmosphere.[53] [54]
8
Radon
9
In 1971, Apollo 15 passed 110 km (68 mi) above the Aristarchus plateau on the Moon, and detected a significant rise in alpha particles thought to be caused by the decay of 222Rn. The presence of 222Rn has been inferred later from data obtained from the Lunar Prospector alpha particle spectrometer.[55] Radon is found in some petroleum. Because radon has a similar pressure and temperature curve as propane, and oil refineries separate petrochemicals based on their boiling points, the piping carrying freshly separated propane in oil refineries can become partially radioactive due to radon decay particles. Residues from the oil and gas industry often contain radium and its daughters. The sulfate scale from an oil well can be radium rich, while the water, oil, and gas from a well often contains radon. The radon decays to form solid radioisotopes which form coatings on the inside of pipework. An oil processing plant, the area of the plant where propane is processed, is often one of the more contaminated areas of the plant as radon has a similar boiling point as propane.[56]
Accumulation in houses Typical domestic exposures are of ≈ 100 Bq/m3 indoors. Depending on how houses are built and ventilated, radon may accumulate in basements and dwellings. Radon concentrations in the same location may differ by a factor of two over a period of 1 hour. Also, the concentration in one room of a building may be significantly different than the concentration in an adjoining room.[1] The distribution of radon concentrations tends to be
Typical Lognormal radon distribution in dwellings.
asymmetrical around the average, the larger concentrations have a disproportionately greater weight. Indoor radon concentration is usually assumed to follow a lognormal distribution on a given territory.[57] Thus, the geometric mean is generally used for estimating the "average" radon concentration in an area.[58] The mean concentration ranges from less than 10 Bq/m3 to over 100 Bq/m3 in some European countries.[59] Typical geometric standard deviations found in studies range between 2 and 3, meaning (given the 68-95-99.7 rule) that the radon concentration is expected to be more than a hundred time the mean concentration for 2 to 3% of the cases. The highest average radon concentrations in the United States are found in Iowa and in the Appalachian Mountain areas in southeastern Pennsylvania.[60] Some of the highest readings ever have been recorded in the Irish town of Mallow, County Cork, prompting local fears regarding lung cancer. Iowa has the highest average radon concentrations in the United States due to significant glaciation that ground the granitic rocks from the Canadian Shield and deposited it as soils making up the rich Iowa farmland.[61] Many cities within the state, such as Iowa City, have passed requirements for radon-resistant construction in new homes. In a few locations, uranium tailings have been used for landfills and were subsequently built on, resulting in possible increased exposure to radon.[1]
Radon
10
Industrial production Radon is obtained as a by-product of uraniferous ores processing after transferring into 1% solutions of hydrochloric or hydrobromic acids. The gas mixture extracted from the solutions contains H2, O2, He, Rn, CO2, H2O and hydrocarbons. The mixture is purified by passing it over copper at 720 °C to remove the H2 and the O2, and then KOH and P2O5 are used to remove the acids and moisture by sorption. Radon is condensed by liquid nitrogen and purified from residue gases by sublimation.[62] Radon commercialization is regulated, but it is available in small quantities for the calibration of 222Rn measurement systems, at a price of almost $6,000 per milliliter of radium solution (which only contains about 15 picograms of actual radon at a given moment).[63] Radon is produced by a solution of radium-226 (half-life of 1 600 years). Radium-226 decays by alpha-particle emission, producing Radon which collects over 3 samples of radium-226 at a rate of about 1 mm /day per gram of radium; equilibrium is quickly achieved and the radon is produced in a steady flow, with an activity equals that of the radium (50 Bq). Gaseous 222Rn (half-life of ~ four days) escapes from the capsule through diffusion.
Concentration scale Bq/m3
Occurrence example 1 Radon concentration at the shores of large oceans is typically 1 Bq/m3. Radon trace concentration above oceans or in Antarctica can be lower than 0.1 Bq/m3. 10 Mean continental concentration in the open air : 10 to 30 Bq/m3. Based on a series of surveys, the global mean indoor radon concentration is estimated to be 39 Bq/m3.
100 Typical indoor domestic exposure. Most countries have adopted a radon concentration of 200–400 Bq/m3 for indoor air as an Action or Reference Level. If testing shows levels less than 4 picocuries radon per liter of air (160 Bq/m3), then no action is necessary. A cumulated exposure of 230 Bq/m3 of radon gas concentration during a period of 1 year corresponds to 1 WLM. 1,000 Very high radon concentrations (>1000 Bq/m3) have been found in countries where houses are built on soils with a high uranium content and/or high permeability of the ground. For levels are 20 picocuries radon per liter of air (800 Bq/m3) or higher, the home owner should consider some type of procedure to decrease indoor radon levels. 10,000 The "Working Level" in uranium mines corresponds to a 7000 Bq/m3 concentration. The concentration in the air at the (unventilated) Gastein Healing Gallery averages 43 kBq/m3 (about [49] 1.2 nCi/l) with maximal value of 160 kBq/m3 (about 4.3 nCi/l). 100,000 About 100,000 Bq/m3 (2,700 pCi/l) was found in Stanley Watras's basement 1,000,000 Expositions reaching 1,000,000 Bq/m3 can be found in unventilated uranium mines.
Radon
Applications Medical Arthritis It has been said that exposure to radon mitigates auto-immune diseases such as arthritis.[64] As a result, in the late 20th century and early 21st century, some "health mines" were established in Basin, Montana which attracted people seeking relief from health problems such as arthritis through limited exposure to radioactive mine water and radon. The practice is controversial because of the "well-documented ill effects of high-dose radiation on the body."[65] Radon has nevertheless been found to induce beneficial long-term effects.[66] Bathing Radioactive water baths have been applied since 1906 in Jáchymov, Czech Republic, but even before radon discovery they were used in Bad Gastein, Austria. Radium-rich springs are also used in traditional Japanese onsen in Misasa, Tottori prefecture. Drinking therapy is applied in Bad Brambach, Germany. Inhalation therapy is carried out in Gasteiner-Heilstollen, Austria, in Kowary, Poland and in Boulder, Montana, United States. In the United States and Europe there are several "radon spas," where people sit for minutes or hours in a high-radon atmosphere in the belief that low doses of radiation will invigorate or energize them.[67] Medical Radiography The radon gas which is used as a cancer treatment in medicine is obtained from the decay of a radium chloride source. In the past, radium and radon have both been used for X-ray medical radiography, but they have fallen out of use as they are radiotoxic alpha radiation emitters which are expensive and have been replaced with iridium-192 and cobalt-60 since they are far better photon sources. Radiotherapy Radon has been produced commercially for use in radiation therapy, but for the most part has been replaced by radionuclides made in accelerators and nuclear reactors. Radon has been used in implantable seeds, made of gold or glass, primarily used to treat cancers. The gold seeds were produced by filling a long tube with radon pumped from a radium source, the tube being then divided into short sections by crimping and cutting. The gold layer keeps the radon within, and filters out the alpha and beta radiations, while allowing the gamma rays to escape (which kill the diseased tissue). The activities might range from 0.05 to 5 millicuries per seed (2 to 200 MBq).[68] The gamma rays are produced by radon and the first short-lived elements of its decay chain (218Po, 214Pb, 214Bi, 214Po). Radon and its first decay products being very short-lived, the seed is left in place. After 12 half-lives (43 days), radon radioactivity is at 1/2000 of its original level. At this stage, the predominant residual activity is due to the radon decay product 210Pb, whose half-life (22.3 year) is 2000 times that or radon (and whose activity is thus 1/2000 or radon's), and its descendants 210Bi and 210Po, totalizing 0.03% of the initial seed activity. In the early part of the 20th century in the USA, gold which was contaminated with 210Pb entered the jewelry industry. This was from gold seeds which had held 222Rn that had been
11
Radon melted down after the radon had decayed.[69] [70] Wearing a contaminated ring could lead to a skin exposition of 10 to 100 rad/day (0.4 to 4 mSv/h)[71]
Scientific Tracking of air masses Radon emanation from the soil varies with soil type and with surface uranium content, so outdoor radon concentrations can be used to track air masses to a limited degree. This fact has been put to use by some atmospheric scientists. Because of radon's rapid loss to air and comparatively rapid decay, radon is used in hydrologic research that studies the interaction between ground water and streams. Any significant Mean concentration of radon in the atmosphere. concentration of radon in a stream is a good indicator that there are local inputs of ground water. Radon is also used in the dating of oil-containing soils because radon has a high affinity for oil-like substances. Geological faults Radon soil-concentration has been used in an experimental way to map buried close-subsurface geological faults because concentrations are generally higher over the faults. Similarly, it has found some limited use in geothermal prospecting. Some researchers have also looked at elevated soil-gas radon concentrations, or rapid changes in soil or groundwater radon concentrations, for earthquake prediction.[72] The theory is that compression around a fault about to rupture could produce radon emission, as if the ground were being squeezed like a sponge. In the 1970s and 1980s, scientific measurements of radon emissions near faults found that earthquakes often occurred with no radon signal, and radon was often detected with no earthquake to follow. It was then dismissed by many as an unreliable indicator.[73] However, as of 2009, it is under investigation as a possible precursor by NASA.[74]
12
Radon Power Radon is a known pollutant emitted from geothermal power stations, though it disperses rapidly, and no radiological hazard has been demonstrated in various investigations. The trend in geothermal plants is to reinject all emissions by pumping deep underground, and this seems likely to ultimately decrease such radon hazards further.
Health incidence of radon exposition Cancer on miners Radon is a common problem encountered during uranium mining, and significant excesses in deaths from lung cancer have been identified in epidemiology studies of uranium miners and other hard rock miners employed in the 1940s and 1950s.[75] [76] [77] The first major studies with radon and health occurred in the context of uranium mining, first in the Joachimsthal region of Bohemia and then in the Relative risk of lung cancer mortality Southwestern United States during the early Cold War. by cumulative exposure to radon decay Because radon is a product of the radioactive decay of products (in WLM) from the combined data from 11 cohorts of underground uranium, underground uranium mines may have high hard rock miners. Though high concentrations of radon. Many uranium miners in the exposures (>50 WLM) cause Four Corners region contracted lung cancer and other statistically significant excess cancers, pathologies as a result of high levels of exposure to the case of small exposures (10 WLM) is inconclusive and appears slightly radon in the mid-1950s. The increased incidence of lung beneficial in this study. cancer was particularly pronounced among Native American and Mormon miners, because those groups normally have low rates of lung cancer.[78] Safety standards requiring expensive ventilation were not widely implemented or policed during this period.[79] In studies of uranium miners, workers exposed to radon levels of 50 to 150 picocuries of radon per liter of air (2000–6000 Bq/m3) for about 10 years have shown an increased frequency of lung cancer.[1] Statistically significant excesses in lung cancer deaths were present after cumulative exposures of less than 50 WLM.[1] There is, however, unexplained heterogeneity in these results (whose confidence interval do not always overlap).[42] The size of the radon-related increase in lung cancer risk varied by more than an order of magnitude between the different studies.[80] Heterogeneities are possibly due to systematic errors in exposure ascertainment, unaccounted for differences in the study populations (genetic, lifestyle, etc.), or confounding mine exposures.[42] There are a number of confounding factors to consider, including exposure to other agents, ethnicity, smoking history, and work experience. The cases reported in these miners cannot be attributed solely to radon or radon daughters but may be due to exposure to silica, to other mine pollutants, to smoking, or to other causes.[1] [81] The majority of miners in the studies are smokers and all inhale dust and other pollutants in mines. Because radon and cigarette smoke both cause lung-cancer, and since the effect of smoking is far above that of radon, it is complicated to disentangle the effects of the 2 kinds of exposure; misinterpreting the smoking habit by a few percent can blur out
13
Radon the radon effect.[82] Since that time, ventilation and other measures have been used to reduce radon levels in most affected mines that continue to operate. In recent years, the average annual exposure of uranium miners has fallen to levels similar to the concentrations inhaled in some homes. This has reduced the risk of occupationally induced cancer from radon, although it still remains an issue both for those who are currently employed in affected mines and for those who have been employed in the past.[80] The power to detect any excess risks in miners nowadays is likely to be small, because the exposures are much smaller than in the early years of mining.[83]
Health risks Radon-222 has been classified by International Agency for Research on Cancer as being carcinogenic to humans.[84] Raised lung cancer rates have been reported from a number of cohort and case-control studies of underground miners exposed to radon and its decay products. There is sufficient evidence for the carcinogenicity of radon and its decay products in humans for such expositions.[85] The primary route of exposure to radon and its progeny is inhalation. Radiation exposure from radon is indirect. The health hazard from radon does not come primarily from radon itself, but rather from the radioactive products formed in the decay of radon.[1] The general effects of radon to the human body are caused by its radioactivity and consequent risk of radiation-induced cancer. Lung cancer is the only observed consequence of high concentration radon exposures; both human and animal studies indicate that the lung and respiratory system are the primary targets of radon daughter-induced toxicity.[1] Radon has a short half-life (4 days) and decays into other solid particulate radium-series radioactive nuclides. Two of these decay products, polonium-218 and 214, present a significant radiologic hazard.[86] If the gas is inhaled, these radioactive particles are inhaled and may attach to the inner lining of the lung. The pattern of their deposition in the respiratory tract is dependent on whether they are attached to particles or not. These particles remain lodged in the lungs, and continue to decay, causing continued exposure by emitting alpha radiations. The radiation decay products can damage cells in the lung tissue,[87] either create free radicals or cause DNA breaks,[86] perhaps causing mutations that sometimes turn cancerous. The risk of lung cancer caused by smoking is much higher than the risk of lung cancer caused by indoor radon. Radiation from radon has been attributed to increase of lung cancer among smokers too. It is generally believed that exposure to radon and cigarette smoking are synergistic; that is, that the combined effect exceeds the sum of their independent effects. This is because the daughters of radon often become attached to smoke and dust particles, and are then able to lodge in the lungs.[88] It is unknown whether radon causes other types of cancer, but recent studies suggest a need for further studies to assess the relationship between radon and leukemia.[89] [90] The effects of radon, if found in food or drinking water, are unknown. Following ingestion of radon dissolved in water, the biological half-life for removal of radon from the body ranges from 30 to 70 minutes. More than 90% of the absorbed radon is eliminated by exhalation within 100 minutes, By 600 minutes, only 1% of the absorbed amount remains in the body.[1]
14
Radon
Effective dose and cancer risks estimations UNSCEAR recommends[91] a reference value of 9 nSv (Bq·h/m3)−1. This means that people living permanently (8760 h/year) in a high concentration of 1000 Bq/m3 receive a dose of 80 mSv/year. Studies of miners exposed to radon and its decay products provide a direct basis for assessing their lung cancer risk. The BEIR VI report, entitled Health Effects of Exposure to Radon,[82] reported an excess relative risk from exposure to radon that was equivalent to 1.8 % per megabecquerel hours per cubic meter (MBq·h/m3) (95% confidence interval: 0.3, 35) for miners with cumulative exposures below 30 MBq·h/m3.[83] Estimates of risk per unit exposure are 5.38×10−4 per WLM ; 9.68×10−4/WLM for ever smokers ; and 1.67×10−4 per WLM for never smokers.[42] According to the UNSCEAR modelization, based on these miner's studies, the excess relative risk from long-term residential exposure to radon at 100 Bq/m3 is considered to be about 0.16 (after correction for uncertainties in exposure assessment), with about a threefold factor of uncertainty higher or lower than that value.[83] In other words, the absence of ill effects (or even positive hormesis effects) at 100 Bq/m3 are compatible with the known data. The ICPR 65 model[92] follows the same approach, and estimates the relative life long risk probability of radon-induced cancer death to 1.23 x 10−6 / (Bq/m3.year)[93] . This relative risk is a global indicator; the risk estimation is independent of sex, age, or smoking habit. Thus, if a smoker's chances of dying of lung cancer are 10 times that of a nonsmoker's, the relative risks for a given radon exposure will be the same according to that model, meaning that the absolute risk of a radon-generated cancer for a smoker is (implicitly) tenfold that of a nonsmoker. The risk estimates correspond to a unit risk of approximately 3–6 × 10−5 per Bq/m3, assuming a lifetime risk of lung cancer of 3%. This means that a person living in an average European house with 50 Bq/m3 has a lifetime excess lung cancer risk of 1.5–3 × 10−3. Similarly, a person living in a house with a high radon concentration of 1000 Bq/m3 has a lifetime excess lung cancer risk of 3–6%, implying a doubling of background lung cancer risk.[94] The BEIR VI model proposed by the National Academy of Sciences of the USA[82] is more complex. It is a multiplicative model that estimates an excess risk per exposure unit. It takes into account age, elapsed time since exposure, and duration and length of exposure, and its parameters allow for taking smoking habits into account.[93] In the absence of other causes of death, the absolute risks of lung cancer by age 75 at usual radon concentrations of 0, 100, and 400 Bq/m3 would be about 0.4%, 0.5%, and 0.7%, respectively, for lifelong nonsmokers, and about 25 times greater (10%, 12%, and 16%) for cigarette smokers.[95] There is great uncertainty in applying risk estimates derived from studies in miners to the effects of residential radon, and direct estimates of the risks of residential radon are needed.[80]
15
Radon
Studies on domestic exposures Major source of natural radiation The largest natural contributor to public radiation dose is radon, a naturally occurring, radioactive gas found in soil and rock[96] , which comprises approximately 55% of the annual background dose. Radon gas levels vary by locality and the composition of the underlying soil and rocks. Radon (at concentrations encountered in mines) was recognized as carcinogenic in the 1980s, in view of the lung cancer statistics for miners' Average radiation doses received in Germany. Radon accounts for half cohorts. Although radon may of the background dose ; and medical doses reach the same levels as present significant risks, background dose. thousands of people annually go to radon-contaminated mines for deliberate exposure to help with the symptoms of arthritis without any serious health effects.[97] The possible danger of radon exposure in dwellings was discovered in 1984 when Stanley Watras, an employee at the Limerick nuclear power plant in Pennsylvania, set off the radiation alarms on his way to work for two weeks while authorities searched for the source of the contamination. They found that the source was high levels of radon—about 100,000 Bq/m3 (2,700 pCi/l)—in his house's basement, and it was not related to the nuclear plant. The risks associated with living in his house were estimated to be equivalent to smoking 135 packs of cigarettes every day. Following this highly publicized event, national radon safety standards were set, and radon detection and ventilation became a standard homeowner concern,[98] though typical domestic expositions are two to three order of magnitude lower (100 Bq/m3, or 2.5 pCi/l). Beginning with the late 1980s, this led to activists forming campaigns to raise awareness of radiation resulting from radon.[99] Radon as a terrestrial source of background radiation is of particular concern because, although on average it is very rare, this intensely radioactive element can be found in high concentrations in many areas of the world. Some of these areas, including Cornwall and Aberdeenshire in the United Kingdom, have high enough natural radiation levels that nuclear licensed sites cannot be built there—the sites would already exceed legal radiation limits before they opened, and the natural topsoil and rock would all have to be disposed of as low-level nuclear waste.[100] People in affected localities can receive up to 10 mSv per year background radiation.[100]
16
Radon This led to a health policy problem: what is the health impact of domestic exposures to radon for the concentrations (100 Bq/m3) typically found in some buildings?
Detection methods When exposure to a carcinogenic substance is suspected, the cause/effect relationship on any given case can never be ascertained. Lung cancer occurs spontaneously, and there is no difference between a "natural" cancer and another one caused by radon (or smoking). Furthermore, it takes years for a cancer to develop, so that determining the past exposure of a case is usually very approximative. The health effect of radon can only be demonstrated through theory and statistical observation. The Study design for epidemiological methods may be of three kinds: • The best proofs are statistical observations of cohorts (predetermined population with known expositions and exhaustive follow-up), like the ones made on miners, or the studies of Hiroshima and Nagasaki survivors. They are efficient, but very costly when the population needs to be a large one. Such studies can only be used when the effect is strong enough, hence, for high exposures. • Alternate proofs are case-control studies (the environment factors of a “case” population is individually determined, and compared to that of a “control″population, to see what the difference might have been, and which factors may be significant), like the ones that have been used to demonstrate the link between lung cancer and smoking. Such studies can identify key factors when the signal/noise ratio is strong enough, but are very sensitive to selection bias, and prone to the existence of confounding factors. • Lastly, ecological studies may be used (where the global environment variables and their global effect on two different populations are compared). Such studies are “cheap and dirty”: they can be easily conducted on very large populations (the whole USA, in Dr Cohen's study), but are prone to the existence of confounding factors, and exposed to the ecological fallacy problem. Furthermore, theory and observation must confirm each other for a relationship to be accepted as fully proven. Even when a statistical link between factor and effect appears significant, it must be backed by a theoretical explanation; and a theory is not accepted as factual unless confirmed by observations.
17
Radon
18
Epidemiology studies of domestic expositions Cohort studies are impractical for the study of domestic radon exposure. The expected effect of small exposures being very small, the direct observation of this effect would require huge cohorts: the populations of whole countries. Several ecological studies have been performed to assess possible relationships between selected cancers and estimated radon levels within particular geographic regions where environmental radon levels appear to be higher than other geographic regions.[103] Results of such ecological studies are mixed; both positive and negative associations, as well as no significant associations, have been suggested.[104] The most direct way to assess the risks posed by radon in homes is through case-control studies. The studies have not produced a definitive answer, primarily because the risk is likely to be very small at the low exposure encountered from most homes and because it is difficult to estimate radon exposures that people have received over their lifetimes. In addition, it is clear that far more lung cancers are caused by smoking than are caused by radon.[82]
A controversial epidemiological study showing increased cancer risk vs. radon domestic exposure (5 pCi/l ≈ [101] 200 Bq/m3). This study lacks individual level controls for smoking and radon exposure, and therefore lacks statistical power to draw conclusions on what would be radiation hormesis. Because of this the error bars (which simply reflect the raw data variability) are probably too [102] small.
Epidemiologic radon studies have found trends to increased lung cancer risk from radon with a no evidence of a threshold, and evidence against a threshold above high as 150 Bq/m3 (almost exactly the EPA's action level of 4 pCi/l).[95] Another study similarly found that there is no evidence of a threshold but lacked the statistical power to clearly identify the threshold at this low level.[105] Notably, the latter deviance from zero at low level convinced the World Health Organization that, "The dose-response relation seems to be linear without evidence of a threshold, meaning that the lung cancer risk increases proportionally with increasing radon exposure."[106] The most elaborate case-control epidemiologic radon study performed by R. William Field and colleagues identified a 50% increased lung cancer risk with prolonged radon exposure at the EPA's action level of 4 pCi/l.[107] Iowa has the highest average radon concentrations in the nation and a very stable population which added to the strength of the study. For that study, the odds ratio was found to be increased slightly above the confidence interval (95% CI) for cumulative radon exposures above 17 WLM (6.2 pC/l=230 Bq/m3 and above). The results of a methodical ten-year-long, case-controlled study of residential radon exposure in Worcester County, Massachusetts, found an apparent 60% reduction in lung cancer risk amongst people exposed to low levels (0–150 Bq/m3) of radon gas; levels typically encountered in 90% of American homes—an apparent support for the idea of radiation hormesis.[108] In that study, a significant result (95% CI) was obtained for the 75-150 Bq/m3 category. The study paid close attention to the cohort's levels of smoking, occupational exposure to carcinogens and education attainment. However, unlike the majority of the residential radon studies, the study was not population-based. Errors in
Radon retrospective exposure assessment could not be ruled out in the finding at low levels. Other studies into the effects of domestic radon exposure have not reported a hormetic effect; including for example the respected "Iowa Radon Lung Cancer Study" of Field et al. (2000), which also used sophisticated radon exposure dosimetry.[107] All these case-control studies should be taken with caution, since it is known that just like ecological studies, they are prone to Ecological fallacy and confounding factors, and cannot cannot prove causation. Such studies can be used to point the way, but further confirmation must come from cohort studies, animal testings, and biological models of the radon effect. To date, most of the studies have had inadequate power to detect a risk on their own, although a weighted average of the published results is indicative of a risk (estimated relative risk at 100 Bq/m3 compared with 0 Bq/m3 1.06, 95% confidence interval (95% CI) 1.01-1.10.[80] .[106]
Health policy on radon public exposure Dose-effect model retained Radon has been recognized as carcinogenic to humans at high concentrations, based on miners studies. The highly publicized case of Stanley Watras showed that radon concentration could reach the levels found in mines, where it is considered a health hazard. This led to awareness of radiation resulting from radon. But what is the real health impact of radon concentration in dwellings? This is an open question. The only dose-effect relationship available are those of miners cohorts (for much higher expositions), exposed to radon. Studies of Hiroshima and Nagasaki survivors are less informative (the exposition to radon is chronic, localized, and the ionizing radiations are alpha rays). Although low-exposed miners experienced exposures comparable to long-term residence in high radon houses, the mean cumulative exposure among miners is approximately 30-fold higher than that associated with long-term residency in a typical home. It can be concluded from miner studies that when the radon exposure in houses compares to that in mines (above 1000 Bq/m3), radon is a proven health hazard ; but in the 1980s very little was known on the dose-effect relationship, both theoretically and statistically. Researches have been made since the 1980s, both on epidemiological studies and in the radiobiology field. In the radiobiology and carcinogenesis studies, progress has been made in understanding the first steps of cancer development, but not to the point of validating a reference dose-effect model. The only certainty gained is that the process is very complex, the resulting dose-effect response being complex, and most probably not a linear one. Biologically based models have also been proposed that could project substantially reduced carcinogenicity at low doses.[42] [109] [110] In the epidemiological field, no definite conclusion has been reached. However, from the evidence now available, a threshold exposure, that is, a level of exposure below which there is no effect of radon, cannot be excluded.[82] Given the radon distribution observed in dwellings, and the dose-effect relationship proposed by a given model, a theoretical number of victims can be calculated, and serve as a basis for public health policies. With the BEIR VI model, the main health impact (nearly 75% of the death toll) is to be found at low radon concentration exposures, because most of the population (about 90%)
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Radon
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lives in the 0-200 Bq/m3 range.[111] Under this modeling, the best policy is obviously to reduce the radon levels of all homes where the radon level is above average, because this leads to a significant decrease of radon exposition on a significant fraction of the population ; but this effect is predicted in the 0-200 Bq/m3 range, where the linear model has its maximum uncertainty. From the statistical evidence available, a threshold exposure cannot be excluded ; if such a threshold exists, the real radon health impact would in fact be limited to those homes where the radon concentrations reaches that observed in mines - at most a few percent. If a radiation hormesis effect exists after all, the situation would be even worse : under that hypothesis, suppressing the natural low exposure to radon (in the 0-200 Bq/m3 range) would actually lead to an increase of cancer incidence, due to the suppression of this (hypothetical) protecting effect. Since the low-dose response is unclear, the choice of a model is very controversial. As the saying goes, “guess if you can, choose if you dare”[112] No conclusive statistics being available for the levels of exposure usually found in homes, the risks posed by domestic exposures is usually estimated on the basis of observed lung-cancer deaths caused by higher exposures in mines, under the assumption that the risk of developing lung-cancer increases linearly as the exposure increases.[82] This was the basis for the model proposed by BEIR IV in the 1980s. The linear no-threshold model has since been kept in a conservative approach by the UNSCEAR[83] report and the BEIR VI and BEIR VII[113] publications, essentially for lack of a better choice: Until the [...] uncertainties on low-dose response are resolved, the Committee believes that [the linear no-threshold model] is consistent with developing knowledge and that it remains, accordingly, the most scientifically defensible approximation of low-dose response. However, a strictly linear dose response should not be expected in all circumstances. The BEIR VI committee adopted the linear no-threshold assumption based on its understanding of the mechanisms of radon-induced lung cancer, but recognized that this understanding is incomplete and that therefore the evidence for this assumption is not conclusive.[42]
Death toll attributed to radon In discussing these figures, it should be kept in mind that both the radon distribution in dwelling and its effect at low exposures are not precisely known, and the radon health impact has to be computed (deaths caused by radon domestic exposure cannot be observed as such). These estimations are strongly dependent on the model retained. According to these models, radon exposure is thought to be the second major cause of lung cancer after smoking.[97] Iowa has the highest average radon concentration in the United States; studies performed there have demonstrated a 50% increased lung cancer risk with prolonged radon exposure above the EPA's action level of 4 pCi/l.[107] [114] Based on studies carried out by the National Academy of Sciences in the United States, radon would thus be the second leading cause of lung cancer after smoking, and accounts for 15,000 to 22,000 cancer deaths per year in the US alone.[115] The United States Environmental Protection Agency (EPA) says that radon is the number one cause of lung cancer among non-smokers.[116] The general population is exposed to small amounts of polonium as a radon daughter in indoor air; the isotopes 214Po and 218Po are thought to cause the majority[117] of the estimated 15,000–22,000 lung cancer deaths in the US every
Radon year that have been attributed to indoor radon.[118] The Surgeon General of the United States has reported that over 20,000 Americans die each year of radon-related lung cancer.[119] In the United Kingdom, residential radon would be, after cigarette smoking, the second most frequent cause of lung cancer deaths: according to models, 83.9% of deaths are attributed to smoking only, 1.0% to radon only, and 5.5% to a combination of radon and smoking.[80]
Radon concentration guidelines The European Union recommends that action should be taken starting from concentrations of 400 Bq/m3 (11 pCi/l) for old houses and 200 Bq/m3 (5 pCi/l) for new ones. After publication of the North American and European Pooling Studies, Health Canada proposed a new guideline that lowers their action level from 800 to 200 Bq/m3 (22 to 5 pCi/l).[120] The United States Environmental Protection Agency (EPA) strongly Predicted fraction of homes exceeding the EPA's recommended recommends action for any house action level of 4 pCi/l with a concentration higher than 3 [87] 148 Bq/m (4 pCi/l), and encourages action starting at 74 Bq/m3 (2 pCi/l). EPA recommends that all homes should be monitored for radon. If testing shows levels less than 4 picocuries radon per liter of air (160 Bq/m3), then no action is necessary. For levels are 20 picocuries radon per liter of air (800 Bq/m3) or higher, the home owner should consider some type of procedure to decrease indoor radon levels.[1] For instance, since radon has a half-life of four days, opening the windows once a day can cut the mean radon concentration to one fourth of its level. The EPA recommends homes be fixed if an occupant's long-term exposure will average 4 picocuries per liter (pCi/l) (148 Bq m−3) or higher.[121] The United States Environmental Protection Agency (EPA) estimates that one in 15 homes in the United States has radon levels above the recommended guideline of 4 picocuries per liter (pCi/l) (148 Bq/m3).[87] EPA radon risk level tables including comparisons to other risks encountered in life are available in their citizen's guide.[122] The EPA estimates that nationally, 8% to 12% of all houses are above their maximum "safe levels" (four picocuries per liter—the equivalent to roughly 200 chest x-rays). The United States Surgeon General and the EPA both recommend that all homes be tested for radon. The limits retained do not correspond to a known threshold in the biological effect, but are determined by a cost-efficiency analysis. EPA believes that a 150 Bq/m3 level (4 pCi/l) is achievable in the vast majority of homes for a reasonable cost, the average cost per life saved by using this action level is about $700,000.[123]
21
Radon
22
For radon concentration in drinkable water, the World Health Organization issued as guidelines (1988) that remedial action should be considered when the radon activity exceeded 100 kBq/m3 in a building, and remedial action should be considered without long delay if exceeding 400 kBq/m3.[1]
Radon testing There are relatively simple tests for radon gas, but these tests are not commonly done, even in areas of known systematic hazards. Radon test kits are commercially available. The short-term radon test kits used for screening purposes are inexpensive, in many cases free. The kit includes a collector that the user hangs in the lowest livable floor of the house for 2 to 7 days. The user then sends the collector to a laboratory for analysis. The
National
Environmental
Health
A radon test kit.
Association
provides a list of radon measurement professionals.[124] Long term kits, taking collections for up to one year, are also available. An open-land test kit can test radon emissions from the land before construction begins. The EPA and the National Environmental Health Association have identified 15 types of radon testing.[125] A Lucas cell is one type of device. Radon levels fluctuate naturally. An initial test might not be an accurate assessment of a home's average radon level. Transient weather can affect short term measurements.[126] Therefore, a high result (over 4 pCi/l) justifies repeating the test before undertaking more expensive abatement projects. Measurements between 4 and 10 pCi/l warrant a long term radon test. Measurements over 10 pCi/l warrant only another short term test so that abatement measures are not unduly delayed. Purchasers of real estate are advised to delay or decline a purchase if the seller has not successfully abated radon to 4 pCi/l or less. Since radon concentrations vary substantially from day to day, single grab-type measurements are generally not very useful, except as a means of identifying a potential problem area, and indicating a need for more sophisticated testing.[1]
Mitigation Transport of radon in indoor air is almost entirely controlled by the ventilation rate in the enclosure. Generally, the indoor radon concentrations increase as ventilation rates decrease.[1] In a well ventilated place, the radon concentration tends to align with outdoor values (typically 10 Bq/m3, ranging from 1 to 100 Bq/m3). Radon levels in indoor air can be lowered in a number of ways, from sealing cracks in floors and walls to increasing the ventilation rate of the building. The five principal ways of reducing the amount of radon accumulating in a house are:[106] • Improving the ventilation of the house and avoiding the transport of radon from the basement into living rooms; • Increasing under-floor ventilation; • Installing a radon sump system in the basement; • Sealing floors and walls; and
Radon • Installing a positive pressurization or positive supply ventilation system. The half-life for radon is 3.8 days, indicating that once the source is removed, the hazard will be greatly reduced within a few weeks. Positive-pressure ventilation systems can be combined with a heat exchanger to recover energy in the process of exchanging air with the outside, and simply exhausting basement air to the outside is not necessarily a viable solution as this can actually draw radon gas into a dwelling. Homes built on a crawl space may benefit from a radon collector installed under a "radon barrier" (a sheet of plastic that covers the crawl space). ASTM E-2121 is a standard for reducing radon in homes as far as practicable below 4 picocuries per liter (pCi/l) in indoor air.[127] [128] The National Environmental Health Association and the National Radon Safety Board administer voluntary National Radon Proficiency Programs for radon professionals consisting of individuals and companies wanting to take training courses and examinations to demonstrate their competency.[129] A list of mitigation service providers is available.[130] Indoor radon can be mitigated by sealing basement foundations, water drainage, or by sub-slab de-pressurization. In severe cases, mitigation can use air pipes and fans to exhaust sub-slab gases to the outside. Indoor ventilation systems are more effective, but exterior ventilation can be cost-effective in some cases.
See also • • • • • •
International Radon Project Lucas cell Radon mitigation Radon removal Radiation Exposure Compensation Act Radiohalo
External links • Toxicological Profile for Radon [131], Draft for Public Comment, Agency for Toxic Substances and Disease Registry, September 2008 • Health Effects of Exposure to Radon: BEIR VI. Committee on Health Risks of Exposure to Radon (BEIR VI), National Research Council available on-line [132] • UNSCEAR 2000 Report to the General Assembly [133],with scientific annexes : Annex B: Exposures from natural radiation sources. • Technical support document for the 1992 Citizen's guide to radon [134], [EPA 400-R-92-011, May 1992]. • Should you measure the radon concentration in your home? [135], Phillip N. Price, Andrew Gelman, in Statistics: A Guide to the Unknown, January 2004.
23
Radon
External links • Radon [136] and radon publications [134] at the United States Environmental Protection Agency • Radon Information Center [137] • Frequently Asked Questions About Radon [138] at National Safety Council • Home Buyer's and Seller's Guide to Radon [139] An article by the International Association of Certified Home Inspectors (InterNACHI) • Radon and Lung Health from the American Lung Association [140] • Radon's impact on your health - Lung Association [141] pnb:نوڈیر
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Radon [21] Le radon, aspects historiques et perception du risque (http:/ / www. radon-france. com/ pdf/ historique. pdf), Roland Masse. [22] Radon Toxicity : Who is at Risk? (http:/ / www. atsdr. cdc. gov/ csem/ radon/ whosat_risk. html), Agency for Toxic Substances and Disease Registry, 2000. [23] " Radon (http:/ / www. allmeasures. com/ Formulae/ d1/ d2/ d3/ Results2_2. asp?formula=2& material=Radon+ [Rn])". All Measures. 2004. . Retrieved 2008-02-12. [24] Williams, David R. (2007-04-19). " Earth Fact Sheet (http:/ / nssdc. gsfc. nasa. gov/ planetary/ factsheet/ earthfact. html)". NASA. . Retrieved 2008-06-26. [25] " Radon (http:/ / education. jlab. org/ itselemental/ ele086. html)". Jefferson Lab. . Retrieved 2008-06-26. [26] Thomas http:/ / books. google. de/ books?hl=de& lr=& id=T0Iiv0BJ1E0C& oi=fnd& pg=PA13, Jens (2002). Noble Gases. Marshall Cavendish. ISBN 9780761414629. [27] Sonzogni, Alejandro. " Interactive Chart of Nuclides (http:/ / www. nndc. bnl. gov/ chart/ )". National Nuclear Data Center: Brookhaven National Laboratory. . Retrieved 2008-06-06. [28] " Principal Decay Scheme of the Uranium Series (http:/ / www. gulflink. osd. mil/ library/ randrep/ du/ mr1018. 7. appa. html)". Gulflink.osd.mil. . Retrieved 2008-09-12. [29] Bader, Richard F.W.. " An Introduction to the Electronic Structure of Atoms and Molecules (http:/ / miranda. chemistry. mcmaster. ca/ esam/ )". McMaster University. . Retrieved 2008-06-26. [30] David R. Lide (2003). "Section 10, Atomic, Molecular, and Optical Physics; Ionization Potentials of Atoms and Atomic Ions". CRC Handbook of Chemistry and Physics, 84th Edition. Boca Raton, Florida: CRC Press. [31] Avrorin, V V (1982). "The Chemistry of Radon". Russian Chemical Reviews 51: 12. doi: 10.1070/RC1982v051n01ABEH002787 (http:/ / dx. doi. org/ 10. 1070/ RC1982v051n01ABEH002787). [32] Stein, L. (1970). "Ionic Radon Solution". Science 168: 362. doi: 10.1126/science.168.3929.362 (http:/ / dx. doi. org/ 10. 1126/ science. 168. 3929. 362). PMID 17809133. [33] Pitzer, Kenneth S. (1975). "Fluorides of radon and element 118". J. Chem. Soc., Chem. Commun.: 760–1. doi: 10.1039/C3975000760b (http:/ / dx. doi. org/ 10. 1039/ C3975000760b). [34] Meng- Sheng Liao; Qian- Er Zhang (1998). "Chemical Bonding in XeF2, XeF4, KrF2, KrF4, RnF2, XeCl2, and XeBr2: From the Gas Phase to the Solid State". The Journal of Physical Chemistry A 102: 10647. doi: 10.1021/jp9825516 (http:/ / dx. doi. org/ 10. 1021/ jp9825516). [35] Filatov, Michael (2003). "Bonding in radon hexafluoride: An unusual relativistic problem?". Physical Chemistry Chemical Physics 5: 1103. doi: 10.1039/b212460m (http:/ / dx. doi. org/ 10. 1039/ b212460m). [36] Holloway, J (1986). "Noble-gas fluorides". Journal of Fluorine Chemistry 33: 149. doi: 10.1016/S0022-1139(00)85275-6 (http:/ / dx. doi. org/ 10. 1016/ S0022-1139(00)85275-6). [37] Avrorin, V. V.; Krasikova, R. N.; Nefedov, V. D.; Toropova, M. A. (1982). "The Chemistry of Radon". Russ. Chem. Review 51: 12. doi: 10.1070/RC1982v051n01ABEH002787 (http:/ / dx. doi. org/ 10. 1070/ RC1982v051n01ABEH002787). [38] Malli, Gulzari L. (2002). "Prediction of the existence of radon carbonyl: RnCO". International Journal of Quantum Chemistry 90: 611. doi: 10.1002/qua.963 (http:/ / dx. doi. org/ 10. 1002/ qua. 963). [39] Runeberg, Nino (1998). "Relativistic pseudopotential calculations on Xe2, RnXe, and Rn2: The van der Waals properties of radon". International Journal of Quantum Chemistry 66: 131. doi: 10.1002/(SICI)1097-461X(1998)66:2<131::AID-QUA4>3.0.CO;2-W (http:/ / dx. doi. org/ 10. 1002/ (SICI)1097-461X(1998)66:2<131::AID-QUA4>3. 0. CO;2-W). [40] Browne, Malcolm W. (1993-03-05). " Chemists Find Way to Make An 'Impossible' Compound - New York Times (http:/ / query. nytimes. com/ gst/ fullpage. html?res=9F0CE2DE1E3CF936A35750C0A965958260& sec=& spon=& pagewanted=all)". Query.nytimes.com. . Retrieved 2009-01-30. [41] " Why Measure RDPs? (http:/ / www. progenygrp. com/ why_measure_rdps. htm)". . Retrieved 2009-07-07. [42] " EPA Assessment of Risks from Radon in Homes (http:/ / www. epa. gov/ radon/ pdfs/ 402-r-03-003. pdf)". Office of Radiation and Indoor Air, US Environmental Protection Agency. June 2003. . [43] CIPR 65 [44] Yamamoto, M. (2006). "Radon". Journal of Environmental Radioactivity 86: 110. doi: 10.1016/j.jenvrad.2005.08.001 (http:/ / dx. doi. org/ 10. 1016/ j. jenvrad. 2005. 08. 001). [45] " HEALTH HAZARD DATA (http:/ / www. us. lindegas. com/ International/ Web/ LG/ US/ MSDS. nsf/ NotesMSDS/ Air+ 002/ $file/ Air+ 002. pdf)" (PDF). The Linde Group. . Retrieved 2008-06-26. [46] " Le Radon. Un gaz radioactif naturel (http:/ / www. laradioactivite. com/ fr/ site/ pages/ radon. htm)". . Retrieved 2009-07-07. [47] French CEA note on Radon (http:/ / www-carmin. cea. fr/ espace-pedagogique/ rayonnements-ionisants-et-sante/ les-radionucleides/ radon-rn) [48] See for instance Sperrin, Malcolm; Gillmore, Gavin; Denman, Tony (2001). "Radon concentration variations in a Mendip cave cluster". Environmental Management and Health 12: 476. doi: 10.1108/09566160110404881 (http:/ / dx. doi. org/ 10. 1108/ 09566160110404881).
25
Radon [49] Zdrojewicz, Zygmunt (2006). "Radon Treatment Controversy, Dose Response". Dose-Response 4: 106. doi: 10.2203/dose-response.05-025.Zdrojewicz (http:/ / dx. doi. org/ 10. 2203/ dose-response. 05-025. Zdrojewicz). [50] Steck, Daniel J.; Field, R. William; Lynch, Charles F. (1999). " Exposure to Atmospheric Radon (http:/ / www. ehponline. org/ members/ 1999/ 107p123-127steck/ steck-full. html)". Environmental Health Perspectives 107 (2): 123. doi: 10.2307/3434368 (http:/ / dx. doi. org/ 10. 2307/ 3434368). . [51] Field, R. William. " Radon Occurrence and Health Risk (http:/ / www. cheec. uiowa. edu/ misc/ radon_occ. pdf)" (PDF). Department of Occupational and Environmental Health, University of Iowa. . Retrieved 2008-02-02. [52] " The Clinical Principles Of Balneology & Physical Medicine (https:/ / www. amtamassage. org/ journal/ winter03_journal/ balneology. html)". . Retrieved 2009-07-07. [53] " The Geology of Radon (http:/ / energy. cr. usgs. gov/ radon/ georadon/ 3. html)". United States Geological Survey. . Retrieved 2008-06-28. [54] " Radon-222 as a tracer in groundwater-surface water interactions (http:/ / www. cosis. net/ abstracts/ EGU2008/ 08953/ EGU2008-A-08953. pdf?PHPSESSID=)" (PDF). Lancaster University. . Retrieved 2008-06-28. [55] Lawson, S.; Feldman, W.; Lawrence, D.; Moore, K.; Elphic, R.; Belian, R. (2005). "Recent outgassing from the lunar surface: the Lunar Prospector alpha particle spectrometer". J. Geophys. Res. 110: 1029. doi: 10.1029/2005JE002433 (http:/ / dx. doi. org/ 10. 1029/ 2005JE002433). [56] " Potential for Elevated Radiation Levels In Propane (http:/ / www. neb-one. gc. ca/ clf-nsi/ rsftyndthnvrnmnt/ sfty/ sftydvsr/ 1994/ nbs199401-eng. pdf)". National Energy Board. April 1994. . Retrieved 2009-07-07. [57] Numerous references, see for instance Analysis And Modelling Of Indoor Radon Distributions Using Extreme Values Theory (http:/ / www. geology. cz/ extranet/ vav/ geochemie-zp/ radon/ sympozia/ 2006/ radon-2006-258-265. pdf) or Indoor Radon in Hungary (Lognormal Mysticism) (http:/ / www. geology. cz/ extranet/ vav/ geochemie-zp/ radon/ sympozia/ 2006/ radon-2006-252-257. pdf) for a discussion. [58] " Data Collection and Statistical Computations (http:/ / aprg. utoledo. edu/ radon/ datacoll. html)". . Retrieved 2009-07-07. [59] " Sources to effects assessment for radon in homes and workplaces (http:/ / www-prositon. cea. fr/ var/ plain/ storage/ original/ media/ UNSCEAR_radon. pdf)". UNSCEAR. . Retrieved 2009-07-07. [60] Price, Phillip N.; Nero, A.; Revzan, K.; Apte, M.; Gelman, A.; Boscardin, W. John. " Predicted County Median Concentration (http:/ / eetd. lbl. gov/ IEP/ high-radon/ USgm. htm)". Lawrence Berkeley National Laboratory. . Retrieved 2008-02-12. [61] Field, R. William. 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Francium
1
Francium radon ← francium → radium Cs ↑ Fr ↓ Uue Periodic Table Extended Periodic Table General Name, symbol, number
francium, Fr, 87
Element category
alkali metals
Group, Period, Block
1, 7, s
Appearance
metallic Standard atomic weight
−1
(223) g·mol
Electron configuration
1
[Rn] 7s
Electrons per shell
2, 8, 18, 32, 18, 8, 1 Physical properties
Phase
solid Density (near r.t.)
1.87 g·cm
Melting point
−3
? 300 K (? 27 °C, ? 80 °F) Boiling point
? 950 K (? 677 °C, ? 1250 °F)
Heat of fusion
ca. 2 kJ·mol−1
Heat of vaporization
ca. 65 kJ·mol−1 Vapor pressure (extrapolated)
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
404
454
519
608
738
946
Atomic properties Crystal structure Oxidation states Electronegativity Ionization energies
body centered cubic 1 (strongly basic oxide) 0.7 (Pauling scale) 1st: 380 kJ/mol
Francium
2 Covalent radius
260 pm Miscellaneous
Magnetic ordering
Paramagnetic
Electrical resistivity
3 µΩ·m
Thermal conductivity
(300 K) 15 W·m−1·K−1
CAS registry number
7440-73-5 Most-stable isotopes
Main article: Isotopes of francium iso 221
Fr
NA syn
half-life
DM
4.8 min
α
217
–
2.033
222
–
1.149
223
α
5.430
219
syn
14.2 min
β
223
trace
22.00 min
β
Fr
DP
6.457
222
Fr
DE (MeV)
At Ra Ra At
References
Francium (pronounced /ˈfrænsiəm/), formerly known as eka-caesium and actinium K,[1] is a chemical element that has the symbol Fr and atomic number 87. It has the lowest electronegativity of all known elements, and is the second rarest naturally occurring element (after astatine). Francium is a highly radioactive metal that decays into astatine, radium, and radon. As an alkali metal, it has one valence electron. Francium was discovered by Marguerite Perey in France (from which the element takes its name) in 1939. It was the last element discovered in nature, rather than synthesized.[2] Outside the laboratory, francium is extremely rare, with trace amounts found in uranium and thorium ores, where the isotope francium-223 continually forms and decays. As little as 20-30 g (one ounce) exists at any given time throughout the Earth's crust; the other isotopes are entirely synthetic. The largest amount ever collected of any isotope was a cluster of 10,000 atoms (of francium-210) created as an ultracold gas at Stony Brook in 1997.[3]
Characteristics Francium is the least stable of the naturally occurring elements: its most stable isotope, francium-223, has a maximum half-life of only 22 minutes. In contrast, astatine, the second-least stable naturally occurring element, has a maximum half-life of 8.5 hours.[4] All isotopes of francium decay into either astatine, radium, or radon.[4] Francium is also less stable than all synthetic elements up to element 105.[3] Francium is an alkali metal whose chemical properties most resemble those of caesium.[3] A very heavy element with a single valence electron,[5] it has the highest equivalent weight of any element.[3] Liquid francium — if such a substance were to be created — should have a surface tension of 0.05092 N/m at its melting point.[6] The francium’s melting point was claimed to have calculated to be around 27 °C (80 °F, 300 K). However, the melting point is
Francium uncertain because of the element’s extreme rarity and radioactivity. This melting point may have been in limited precision, or so much heat produced from radioactivity that its calculated melting point may have been overestimated.[7] However, the melting point of francium is estimated to be about 22 °C (71 °F, 295 K), based from the periodic trends in melting points with other alkali metals. Also the boiling point may have been overestimated at around 677 °C (1250 °F, 950 K). Based from the periodic trends with other alkali metals, the boiling point of francium is estimated to be between 660 to 665 °C (1220 to 1230 °F, 935 to 940 K). Because radioactive elements give off heat Francium would almost certainly be a liquid if enough visible Francium were to be produced. Linus Pauling estimated the electronegativity of francium at 0.7 on the Pauling scale, the same as caesium;[8] the value for caesium has since been refined to 0.79, although there are no experimental data to allow a refinement of the value for francium.[9] Francium has a slightly higher ionisation energy than caesium,[10] 392.811(4) kJ/mol as opposed to 375.7041(2) kJ/mol for caesium, as would be expected from relativistic effects, and this would imply that caesium is the less electronegative of the two. Francium coprecipitates with several caesium salts, such as caesium perchlorate, which results in small amounts of francium perchlorate. This coprecipitation can be used to isolate francium, by adapting the radiocaesium coprecipitation method of Glendenin and Nelson. It will additionally coprecipitate with many other caesium salts, including the iodate, the picrate, the tartrate (also rubidium tartrate), the chloroplatinate, and the silicotungstate. It also coprecipitates with silicotungstic acid, and with perchloric acid, without another alkali metal as a carrier, which provides other methods of separation.[11] [12] Nearly all francium salts are water-soluble.[13]
Applications Due to its instability and rarity, there are no commercial applications for francium.[14] [15] [16] [17] [18] It has been used for research purposes in the fields of biology and of atomic structure. Its use as a potential diagnostic aid for various cancers has also been explored,[4] but this application has been deemed impractical.[16] Francium's ability to be synthesized, trapped, and cooled, along with its relatively simple atomic structure have made it the subject of specialized spectroscopy experiments. These experiments have led to more specific information regarding energy levels and the coupling constants between subatomic particles.[19] Studies on the light emitted by laser-trapped francium-210 ions have provided accurate data on transitions between atomic energy levels which are fairly similar to those predicted by quantum theory.[20]
3
Francium
History As early as 1870, chemists thought that there should be an alkali metal beyond caesium, with an atomic number of 87.[4] It was then referred to by the provisional name eka-caesium.[21] Research teams attempted to locate and isolate this missing element, and at least four false claims were made that the element had been found before an authentic discovery was made.
Erroneous and incomplete discoveries Russian chemist D. K. Dobroserdov was the first scientist to claim to have found eka-caesium, or francium. In 1925, he observed weak radioactivity in a sample of potassium, another alkali metal, and concluded that eka-caesium was contaminating the sample.[22] He then published a thesis on his predictions of the properties of eka-caesium, in which he named the element russium after his home country.[23] Shortly thereafter, Dobroserdov began to focus on his teaching career at the Polytechnic Institute of Odessa, and he did not pursue the element further.[22] The following year, English chemists Gerald J. F. Druce and Frederick H. Loring analyzed X-ray photographs of manganese(II) sulfate.[23] They observed spectral lines which they presumed to be of eka-caesium. They announced their discovery of element 87 and proposed the name alkalinium, as it would be the heaviest alkali metal.[22] In 1930, Fred Allison of the Alabama Polytechnic Institute claimed to have discovered element 87 when analyzing pollucite and lepidolite using his magneto-optical machine. Allison requested that it be named virginium after his home state of Virginia, along with the symbols Vi and Vm.[23] [24] In 1934, however, H.G. MacPherson of UC Berkeley disproved the effectiveness of Allison's device and the validity of this false discovery.[25] In 1936, Romanian chemist Horia Hulubei and his French colleague Yvette Cauchois also analyzed pollucite, this time using their high-resolution X-ray apparatus.[22] They observed several weak emission lines, which they presumed to be those of element 87. Hulubei and Cauchois reported their discovery and proposed the name moldavium, along with the symbol Ml, after Moldavia, the Romanian province where they conducted their work.[23] In 1937, Hulubei's work was criticized by American physicist F. H. Hirsh Jr., who rejected Hulubei's research methods. Hirsh was certain that eka-caesium would not be found in nature, and that Hulubei had instead observed mercury or bismuth X-ray lines. Hulubei, however, insisted that his X-ray apparatus and methods were too accurate to make such a mistake. Because of this, Jean Baptiste Perrin, Nobel Prize winner and Hulubei's mentor, endorsed moldavium as the true eka-caesium over Marguerite Perey's recently discovered francium. Perey, however, continuously criticized Hulubei's work until she was credited as the sole discoverer of element 87.[22]
Perey's analysis Eka-caesium was discovered in 1939 by Marguerite Perey of the Curie Institute in Paris, France when she purified a sample of actinium-227 which had been reported to have a decay energy of 220 keV. However, Perey noticed decay particles with an energy level below 80 keV. Perey thought this decay activity might have been caused by a previously unidentified decay product, one which was separated during purification, but emerged again out of the pure actinium-227. Various tests eliminated the possibility of the unknown element being thorium, radium, lead, bismuth, or thallium. The new product exhibited
4
Francium
5
chemical properties of an alkali metal (such as coprecipitating with caesium salts), which led Perey to believe that it was element 87, caused by the alpha decay of actinium-227.[21] Perey then attempted to determine the proportion of beta decay to alpha decay in actinium-227. Her first test put the alpha branching at 0.6%, a figure which she later [26] revised to 1%. Perey named the new isotope actinium-K (now referred to as francium-223)[21] and in 1946, she proposed the name catium for her newly discovered element, as she believed it to be the most electropositive cation of the elements. Irène Joliot-Curie, one of Perey's supervisors, opposed the name due to its connotation of cat rather than cation.[21] Perey then suggested francium, after France. This name was officially adopted by the International Union of Pure and Applied Chemistry in 1949,[4] becoming the second element after gallium to be named after France. It was assigned the symbol Fa, but this abbreviation was revised to the current Fr shortly thereafter.[27] Francium was the last element discovered in nature, rather than synthesized, following rhenium in 1925.[21] Further research into francium's structure was carried out by, among others, Sylvain Lieberman and his team at CERN in the 1970s and 1980s.[28]
Occurrence Natural Francium-223 is the result of the alpha decay of actinium-227 and can be found in trace amounts in uranium and thorium minerals.[3] In a given sample of uranium, there is estimated to be only one francium atom for every 1×1018 uranium atoms.[16] It is also calculated that there is at most 30 g of francium in the earth's crust at any time.[29] This makes it the second rarest element in the crust after astatine.[4] [16]
This sample of uraninite contains about 100,000 atoms (3.3 × 10−20 g) of [16] francium-223 at any given time.
Francium
Synthesized Francium can be synthesized in the nuclear reaction 197Au + 18 O → 210Fr + 5n. This process, developed by Stony Brook Physics, yields francium isotopes with masses of 209, 210, and 211,[31] which are then isolated by the magneto-optic trap (MOT).[30] The production rate of a particular isotope depends on the energy of the oxygen beam. An 18O beam from the Stony Brook LINAC creates 210Fr in the gold target with the nuclear reaction 197Au + 18O = 210Fr + 5n. The production required some time to develop and understand. It was critical In the MOT, a magnetic field is created by the copper to operate the gold target very close to its melting point and to solenoids. Neutral francium make sure that its surface was very clean. The nuclear reaction atoms enter the glass bulb imbeds the francium atoms deep in the gold target, and they from the left and are trapped [30] must be removed efficiently. The atoms diffuse fast to the by lasers. surface of the gold target and are released as ions. The francium ions are guided by electrostatic lenses until they land into a surface of hot yttrium and become neutral again. The francium is then injected into a glass bulb. A magnetic field and retroreflected laser beams cool and confine the atoms. Although the atoms remain in the trap for only about 20 seconds before escaping (or decaying), a steady stream of fresh atoms replaces those lost, keeping the number of trapped atoms roughly constant for minutes or longer. Initially, about 1000 francium atoms were trapped in the experiment. This was gradually improved and is capable of trapping over 300,000 neutral atoms of francium a time. Although these are neutral "metallic" atoms ("francium metal"), they in a gaseous unconsolidated state. Enough francium is trapped that a video camera can capture the light given off by the atoms as they fluoresce. The atoms appear as a glowing sphere about 1 millimeter in diameter. This was the very first time that anyone had ever seen francium. The researchers can now make extremely sensitive measurements of the light emitted and absorbed by the trapped atoms, providing the first experimental results on various transitions between atomic energy levels in francium. Initial measurements show very good agreement between experimental values and calculations based on quantum theory. Other synthesis methods include bombarding radium with neutrons, and bombarding thorium with protons, deuterons, or helium ions.[26] Francium has not yet, as of 2009[32], been synthesized in amounts large enough to weigh.[3] [4] [16] [33]
Isotopes There are 34 known isotopes of francium ranging in atomic mass from 199 to 232.[3] Francium has seven metastable nuclear isomers.[3] Francium-223 and francium-221 are the only isotopes that occur in nature, though the former is far more common.[34] Francium-223 is the most stable isotope with a half-life of 21.8 minutes,[3] and it is highly unlikely that an isotope of francium with a longer half-life will ever be discovered or synthesized.[26] Francium-223 is the fifth product of the actinium decay series as the daughter isotope of actinium-227.[18] Francium-223 then decays into radium-223 by beta decay (1149 keV decay energy), with a minor (0.006%) alpha decay path to astatine-219 (5.4 MeV decay energy).[35]
6
Francium Francium-221 has a half-life of 4.8 minutes.[3] It is the ninth product of the neptunium decay series as a daughter isotope of actinium-225.[18] Francium-221 then decays into astatine-217 by alpha decay (6.457 MeV decay energy).[3] The least stable ground state isotope is francium-215, with a half-life of 0.12 μs. (9.54 MeV alpha decay to astatine-211):[3] Its metastable isomer, francium-215m, is less stable still, with a half-life of only 3.5 ns.[36]
Notes [1] [2] [3] [4]
Actually the least unstable isotope, Fr-223 Some synthetic elements, like technetium, have later been found in nature. CRC Handbook of Chemistry and Physics, 4, CRC, 2006, pp. 12, 0-8493-0474-1 Price, Andy (2004-12-20). " Francium (http:/ / www. andyscouse. com/ pages/ francium. htm)". . Retrieved 2007-03-25. [5] Winter, Mark. " Electron Configuration (http:/ / www. webelements. com/ webelements/ elements/ text/ Fr/ eneg. html)". Francium. The University of Sheffield. . Retrieved 2007-04-18. [6] Kozhitov, L. V.; Kol'tsov, V. B.; Kol'tsov, A. V. (2003-02-21). " Evaluation of the Surface Tension of Liquid Francium (http:/ / search. ebscohost. com/ login. aspx?direct=true& db=aqh& AN=16822434& site=ehost-live)". Inorganic Materials (Springer Science & Business Media B.V.) 39 (11): 1138–1141. doi: 10.1023/A:1027389223381 (http:/ / dx. doi. org/ 10. 1023/ A:1027389223381). . Retrieved 2007-04-14. [7] " What are the characteristics and properties of francium? (http:/ / answers. yahoo. com/ question/ index?qid=20080927201134AAcNUDO)". Yahoo! Answers. 2008. . Retrieved 2009-01-04. [8] Pauling, Linus (1960). The Nature of the Chemical Bond (3rd Edn.). Cornell University Press. pp. 93. [9] Allred, A. L. (1961). "Electronegativity values from thermochemical data". J. Inorg. Nucl. Chem. 17 (3–4): 215–221. doi: 10.1016/0022-1902(61)80142-5 (http:/ / dx. doi. org/ 10. 1016/ 0022-1902(61)80142-5). [10] Andreev, S.V.; Letokhov, V.S.; Mishin, V.I., (1987). " Laser resonance photoionization spectroscopy of Rydberg levels in Fr (http:/ / link. aps. org/ abstract/ PRL/ v59/ p1274)". Phys. Rev. Lett. 59: 1274–76. doi: 10.1103/PhysRevLett.59.1274 (http:/ / dx. doi. org/ 10. 1103/ PhysRevLett. 59. 1274). . [11] Hyde, E. K. (1952). "Radiochemical Methods for the Isolation of Element 87 (Francium)". J. Am. Chem. Soc. 74 (16): 4181–4184. doi: 10.1021/ja01136a066 (http:/ / dx. doi. org/ 10. 1021/ ja01136a066). [12] E. N K. Hyde Radiochemistry of Francium,Subcommittee on Radiochemistry, National Academy of Sciences-National Research Council; available from the Office of Technical Services, Dept. of Commerce, 1960. [13] A. G. Maddock. Radioactivity of the heavy elements. Q. Rev., Chem. Soc., 1951, 3, 270–314. doi: 10.1039/QR9510500270 (http:/ / dx. doi. org/ 10. 1039/ QR9510500270) [14] Winter, Mark. " Uses (http:/ / www. webelements. com/ webelements/ elements/ text/ Fr/ uses. html)". Francium. The University of Sheffield. . Retrieved 2007-03-25. [15] Bentor, Yinon. " Chemical Element.com - Francium (http:/ / www. chemicalelements. com/ elements/ fr. html)". . Retrieved 2007-03-25. [16] Emsley, John (2001). Nature's Building Blocks. Oxford: Oxford University Press. pp. 151–153. ISBN 0-19-850341-5. [17] Gagnon, Steve. " Francium (http:/ / education. jlab. org/ itselemental/ ele087. html)". Jefferson Science Associates, LLC. . Retrieved 2007-04-01. [18] Considine, Glenn D., ed. (2005), "Chemical Elements", Van Nostrand's Encyclopedia of Chemistry, New York: Wylie-Interscience, pp. 332, ISBN 0-471-61525-0 [19] Gomez, E; Orozco, L A, and Sprouse, G D (2005-11-07). " Spectroscopy with trapped francium: advances and perspectives for weak interaction studies (http:/ / www. iop. org/ EJ/ abstract/ 0034-4885/ 69/ 1/ R02/ )". Rep. Prog. Phys. 69 (1): 79–118. doi: 10.1088/0034-4885/69/1/R02 (http:/ / dx. doi. org/ 10. 1088/ 0034-4885/ 69/ 1/ R02). . Retrieved 2007-04-11. [20] Peterson, I (1996-05-11). " Creating, cooling, trapping francium atoms (http:/ / search. ebscohost. com/ login. aspx?direct=true& db=ulh& AN=9605167788& site=src-live)". Science News. pp. 294. . Retrieved 2007-04-11. [21] Adloff, Jean-Pierre; Kaufman, George B. (2005-09-25). Francium (Atomic Number 87), the Last Discovered Natural Element (http:/ / chemeducator. org/ sbibs/ s0010005/ spapers/ 1050387gk. htm). The Chemical Educator 10 (5). Retrieved on 2007-03-26. [22] Fontani, Marco (2005-09-10). " The Twilight of the Naturally-Occurring Elements: Moldavium (Ml), Sequanium (Sq) and Dor (Do) (http:/ / 5ichc-portugal. ulusofona. pt/ uploads/ PaperLong-MarcoFontani. doc)". International Conference on the History of Chemistry. Lisbon. pp. 1–8.
7
Francium [23] Van der Krogt, Peter (2006-01-10). " Francium (http:/ / www. vanderkroft. net/ elements/ elem/ fr. html)". Elementymology & Elements Multidict. . Retrieved 2007-04-08. [24] " Alabamine & Virginium (http:/ / www. time. com/ time/ magazine/ article/ 0,9171,743159,00. html)". TIME. 1932-02-15. . Retrieved 2007-04-01. [25] MacPherson, H. G. (1934-12-21). " An Investigation of the Magneto-Optic Method of Chemical Analysis (http:/ / prola. aps. org/ abstract/ PR/ v47/ i4/ p310_1)". Physical Review (American Physical Society) 47 (4): 310–315. doi: 10.1103/PhysRev.47.310 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 47. 310). . Retrieved 2007-04-08. [26] "Francium", McGraw-Hill Encyclopedia of Science & Technology, 7, McGraw-Hill Professional, 2002, pp. 493–494, ISBN 0-07-913665-6 [27] Grant, Julius (1969), "Francium", Hackh's Chemical Dictionary, McGraw-Hill, pp. 279–280 [28] " History (http:/ / fr. physics. sunysb. edu/ francium_news/ history. HTM)". Francium. SUNY Stony Brook Physics & Astronomy. 2007-02-20. . Retrieved 2007-03-26. [29] Winter, Mark. " Geological information (http:/ / www. webelements. com/ webelements/ elements/ text/ Fr/ geol. html)". Francium. The University of Sheffield. . Retrieved 2007-03-26. [30] " Cooling and Trapping (http:/ / fr. physics. sunysb. edu/ francium_news/ trapping. HTM)". Francium. SUNY Stony Brook Physics & Astronomy. 2007-02-20. . Retrieved 2007-05-01. [31] " Production of Francium (http:/ / fr. physics. sunysb. edu/ francium_news/ production. HTM)". Francium. SUNY Stony Brook Physics & Astronomy. 2007-02-20. . Retrieved 2007-03-26. [32] http:/ / en. wikipedia. org/ wiki/ Francium [33] " Francium (http:/ / periodic. lanl. gov/ elements/ 87. html)". Los Alamos National Laboratory (http:/ / www. lanl. gov/ ) Chemistry Division. 2003-12-15. . Retrieved 2007-03-29. [34] Considine, Glenn D., ed. (2005), "Francium", Van Nostrand's Encyclopedia of Chemistry, New York: Wylie-Interscience, pp. 679, ISBN 0-471-61525-0 [35] National Nuclear Data Center (1990). " Table of Isotopes decay data (http:/ / ie. lbl. gov/ toi/ nuclide. asp?iZA=870223)". Brookhaven National Laboratory. . Retrieved 2007-04-04. [36] National Nuclear Data Center (2003). " Fr Isotopes (http:/ / ie. lbl. gov/ education/ parent/ Fr_iso. htm)". Brookhaven National Laboratory. . Retrieved 2007-04-04.
External links • WebElements.com - Francium (http:/ / www. webelements. com/ webelements/ elements/ text/ Fr/ index. html) • Los Alamos National Laboratory - Francium (http:/ / periodic. lanl. gov/ elements/ 87. html) • Stony Brook University Physics Dept. (http:/ / fr. physics. sunysb. edu/ francium_news/ frconten. htm) . pnb:میسنارف
8
Article Sources and Contributors
Article Sources and Contributors Francium Source: http://en.wikipedia.org/w/index.php?oldid=308667001 Contributors: -Ozone-, 123abcarchy1, 2over0, 65.68.87.xxx, A. di M., ACupOfCoffee, AbJ32, Abrech, Ahoerstemeier, Aitias, Ajedrez, Alexfusco5, AlimanRuna, Alsandro, Amicon, Amire80, Andplus, Andre Engels, AndreasJS, Andres, Andyscouse, Angelofdeath275, Animum, Antandrus, Art LaPella, Atreyu1075, Baccyak4H, Badgersaresexy4, Bammybammy, Bassbonerocks, Beetstra, Binglechild, Blakersquakers, Blehfu, BlueEarth, Bobo192, Boccobrock, Bongwarrior, Brandon5485, Brian0918, Brighterorange, Brossow, Bryan Derksen, CYD, Cacophony, Carnildo, Casliber, Casualtie, Cefalufrnk, Chad okere, Chairman S., Charliehesk, Chessgrand, Chickensunited, Ck lostsword, Colbuckshot, Conversion script, CoolGuy, Cremepuff222, Cryptic C62, Curious Blue, DMacks, DMeyering, Dajwilkinson, Danny, DarkFalls, Darthchaos, Davewild, David Latapie, Dead3y3, Derek Ross, Derek.cashman, Discospinster, DissidentSA, Dohers999, Doulos Christos, Download, Dr. Blofeld, DrBob, DragonflySixtyseven, Dreamafter, Dreish, Dwmyers, Eastlaw, Edgar181, Eeekster, El C, Emperorbma, Eric119, Erik Zachte, Esrever, Ethocmub, Exander2009, Femto, FisherQueen, Focus the moon, Fonzy, Fys, G man yo, Gazimoff, Gene Nygaard, George63, Georgethedecider, Gilliam, HairyPerry, Hak-kâ-ngìn, Hal peridol, HalfShadow, Hall Monitor, Harryboyles, Higgyo, Huntthetroll, Hurley342, IRKAIN, Icairns, IceUnshattered, Ideyal, Igoldste, Indon, Itchweeed, Itub, J.delanoy, JHunterJ, Jaguar2k, Jaraalbe, Jaxl, Jewb6, Jimp, Jklin, Joelholdsworth, John, JohnCD, JoshHood, Jrugordon, Juneappal, Katalaveno, Keith D, Kilo-Lima, Kimchi.sg, KnowledgeOfSelf, Kosebamse, Kozuch, Kukini, Kuru, Kurykh, Kwamikagami, LA2, Ladsgroup, Lightmouse, Looxix, Lychosis, Lysdexia, MER-C, Malcolm, Malcolma, Marc Venot, Markussep, Materialscientist, Mav, Micga, Michaelbusch, Mike Rosoft, Misza13, Morton1991, Morton91, Mschel, Nakon, Nergaal, NewEnglandYankee, Nihil novi, Nihiltres, Nishantsoccer, Nn123645, Nonagonal Spider, Nosorryjustno, O Graeme Burns, Odie5533, Opelio, Oreo Priest, Ossmann, Oxymoron83, Paul-L, Peruvianllama, Phatom87, Philip Trueman, Physchim62, PhysicsIsh, Pinkadelica, Pmanderson, Pol098, Poolkris, Possum, Quadell, Quantpole, Quietust, Rabidmarmcet, Rangek, RedRollerskate, Remember, Rigadoun, Rjwilmsi, Rmrfstar, Rob-bob7-0, RobertG, Robertgreer, Rocket71048576, Roentgenium111, Romanm, Salty!, Salvadoradi, Sam8uk, Saperaud, Schlarpi, Schneelocke, Scog, ScoutingforGirls, Sdavidr2008, Seguin5, Sengkang, Shanel, Shanes, Sl, Slothman32, Smith609, Soliloquial, Squids and Chips, Srtxg, Stebbins, Stifynsemons, Stillwaterslalomforever, SunDragon34, Suruena, Syd Henderson, Szwejk, TDM89034, Tagishsimon, Tarotcards, Tavilis, Tavix, The Librarian, Thingg, Tiki2099, Tombomp, Tpbradbury, Transcendence, Twaz, Until It Sleeps, V1adis1av, VASANTH S.N., Vatic7, Velvetron, Vsmith, Watch37264, Weasrat, WilliamHowardq, WillowW, Wimvandorst, Wknight94, Worm That Turned, X201, Yamamoto Ichiro, Yekrats, Zzuuzz, 436 anonymous edits
Image Sources, Licenses and Contributors image:Fr-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Fr-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Conscious, Paddy, Saperaud Image:Pichblende.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Pichblende.jpg License: Creative Commons Attribution-Sharealike 2.5 Contributors: Original uploader was Kgrr at en.wikipedia Image:franciumtrap.PNG Source: http://en.wikipedia.org/w/index.php?title=File:Franciumtrap.PNG License: unknown Contributors: Cryptic C62
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9
Radium
1
Radium francium ← radium → actinium Ba ↑ Ra ↓ Ubn Periodic Table Extended Periodic Table General Name, symbol, number
radium, Ra, 88
Element category
alkaline earth metals
Group, Period, Block
2, 7, s
Appearance
silvery white metallic Standard atomic weight
−1
(226) g·mol
Electron configuration
2
[Rn] 7s
Electrons per shell
2, 8, 18, 32, 18, 8, 2 Physical properties
Phase
solid Density (near r.t.)
−3
5.5 g·cm
Melting point
973 K (700 °C, 1292 °F) Boiling point
2010 K (1737 °C, 3159 °F)
Heat of fusion
8.5 kJ·mol−1
Heat of vaporization
113 kJ·mol−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
819
906
1037
1209
1446
1799
Atomic properties Crystal structure Oxidation states Electronegativity
body centered cubic 2 (strongly basic oxide) 0.9 (Pauling scale)
Radium
2
Ionization energies
1st: 509.3 kJ/mol 2nd: 979.0 kJ/mol
Covalent radius
221±2 pm Miscellaneous
Magnetic ordering
nonmagnetic
Electrical resistivity
(20 °C) 1 µ Ω·m
Thermal conductivity
(300 K) 18.6 W·m−1·K−1
CAS registry number
7440-14-4 Most-stable isotopes
Main article: Isotopes of radium iso
NA
half-life
DM
DE (MeV)
DP
223
?
11.43 d
alpha
5.99
219
224
?
3.6319 d
alpha
5.789
220
1602 y
alpha
4.871
222
5.75 y
beta
−
0.046
228
Ra Ra
226
Ra
228
Ra
trace ?
Rn Rn Rn Ac
References
Radium (pronounced /ˈreɪdiəm/) is a radioactive chemical element which has the symbol Ra and atomic number 88. Its appearance is almost pure white, but it readily oxidizes on exposure to air, turning black. Radium is an alkaline earth metal that is found in trace amounts in uranium ores. It is extremely radioactive. Its most stable isotope, 226Ra, has a half-life of 1602 years and decays into radon gas.
Characteristics The heaviest of the alkaline earth metals, radium is intensely radioactive and resembles barium in its chemical behavior. This metal is found in tiny quantities in the uranium ore pitchblende, and various other uranium minerals. Radium preparations are remarkable for maintaining themselves at a higher temperature than their surroundings, and for their radiations, which are of three kinds: alpha particles, beta particles, and gamma rays. When freshly prepared, pure radium metal is brilliant white, but blackens when exposed to air (probably due to nitride formation). Radium is luminescent (giving a faint blue color), reacts violently with water and oil to form radium hydroxide and is slightly more volatile than barium. The normal phase of radium is a solid.
Radium
Applications Some of the few practical uses of radium are derived from its radioactive properties. More recently discovered radioisotopes, such as 60Co and 137Cs, are replacing radium in even these limited uses because several of these isotopes are more powerful emitters, safer to handle, and available in more concentrated form. When mixed with beryllium it is a neutron source for physics experiments.
Historical uses Radium was formerly used in self-luminous paints for watches, nuclear panels, aircraft switches, clocks, and instrument dials. More than 100 former watch dial painters who used their lips to shape the paintbrush died from the radiation from the radium that had become stored in their bones. Soon afterward, the adverse effects of radioactivity became widely known. Radium was still used in dials as late as the 1950s. Although tritium's beta radiation is potentially dangerous if ingested, it has replaced radium in these applications. During the 1930s it was found that workers' exposure to radium by handling luminescent paints caused serious health effects which included sores, anemia and bone cancer. This use of radium was stopped soon afterward. This is because radium is treated as calcium by the body, and deposited in the bones, where radioactivity degrades marrow and can mutate bone cells. The litigation and ultimate deaths of five "Radium Girl" employees who had used radium-based luminous paints on the dials of watches and clocks had a significant impact on the formulation of occupational disease labor law. [1] Radium was also put in some foods for taste and as a preservative, but also exposed many people to radiation. Radium was once an additive in products like toothpaste, hair creams, and even food items due to its supposed curative powers.[2] Such products soon fell out of vogue and were prohibited by authorities in many countries, after it was discovered they could have serious adverse health effects. (See for instance Radithor.) Spas featuring radium-rich water are still occasionally touted as beneficial, such as those in Misasa, Tottori, Japan. In the U.S., nasal radium irradiation was also administered to children to prevent middle ear problems or enlarged tonsils from the late 1940s through early 1970s. [3]
In 1909, the famous Rutherford experiment used radium as an alpha source to probe the atomic structure of gold. This experiment led to the Rutherford model of the atom and revolutionised the field of nuclear physics. Radium (usually in the form of radium chloride) was used in medicine to produce radon gas which in turn is used as a cancer treatment, for example several of these radon sources were used in Canada in the 1920s and 1930s.[4] The isotope 223Ra is currently under investigation for use in medicine as cancer treatment of bone metastasis.
History Radium (Latin radius, ray) was discovered by Marie Skłodowska-Curie and her husband Pierre in 1898 in pitchblende coming from North Bohemia, in the Czech Republic (area around Jáchymov). While studying pitchblende the Curies removed uranium from it and found that the remaining material was still radioactive. They then separated out a radioactive mixture consisting mostly of barium which gave a brilliant green flame color and crimson carmine spectral lines which had never been documented before. The Curies
3
Radium
4
announced their discovery to the French Academy of Sciences on 26 December 1898.[5] In 1910, radium was isolated as a pure metal by Curie and André-Louis Debierne through the electrolysis of a pure radium chloride solution by using a mercury cathode and distilling in an atmosphere of hydrogen gas.[6] Radium was first industrially produced in the beginning of the 20th Century by Biraco, a subsidiary company of Union Minière du Haut Katanga (UMHK) in its Olen plant in Belgium. UMHK offered to Marie Curie her first gramme of radium. Historically the decay products of radium were known as radium A, B, C, etc. These are now known to be isotopes of other elements as follows: Isotope Radium emanation
222
Radium A
218
Radium B
214
Radium C
214
Radium C1
214
Radium C2
210
Radium D
210
Radium E
210
Radium F
210
Rn Po Pb Bi Po Tl Pb Bi Po
On February 4, 1936 radium E became the first radioactive element to be made synthetically.[7] One unit for radioactivity, the non-SI curie, is based on the radioactivity of Radioactivity).
226
Ra (see
Occurrence Radium is a decay product of uranium and is therefore found in all uranium-bearing ores. (One ton of pitchblende yields one seventh of a gram of radium).[8] Radium was originally acquired from pitchblende ore from Joachimsthal, Bohemia, in the Czech Republic. Carnotite sands in Colorado provide some of the element, but richer ores are found in the Democratic Republic of the Congo and the Great Lakes area of Canada, and can also be extracted from uranium processing waste. Large radium-containing uranium deposits are located in Canada (Ontario), the United States (New Mexico, Utah, and Virginia), Australia, and in other places.
Radium
5
Compounds Its compounds color flames crimson carmine (rich red or crimson color with a shade of purple) and give a characteristic spectrum. Due to its geologically short half life and intense radioactivity, radium compounds are quite rare, occurring almost exclusively in uranium ores. • • • • • •
radium radium radium radium radium radium
fluoride (RaF2) chloride (RaCl2) bromide (RaBr2) iodide (RaI2) oxide (RaO) nitride (Ra3N2)
Isotopes Radium (Ra) has 25 different known isotopes, four of which are found in nature, with 226Ra being the most common. 223Ra, 224Ra, 226Ra and 228Ra are all generated naturally in the decay of either Uranium (U) or Thorium (Th). 226Ra is a product of 238U decay, and is the longest-lived isotope of radium with a half-life of 1602 years; next longest is 228Ra, a product of 232Th breakdown, with a half-life of 5.75 years.[9]
Radioactivity Radium is over one million times more radioactive than the same mass of uranium. Its decay occurs in at least seven stages; the successive main products have been studied and were called radium emanation or exradio (now identified as radon), radium A (polonium), radium B (lead), radium C (bismuth), etc. Radon is a heavy gas and the later products are solids. These products are themselves radioactive elements, each with an atomic weight a little lower than its predecessor. Radium loses about 1% of its activity in 25 years, being transformed into elements of lower atomic weight with lead being the final product of disintegration. The SI unit of radioactivity is the becquerel (Bq), equal to one disintegration per second. The Curie is a non-SI unit defined as that amount of radioactivity which has the same disintegration rate as 1 gram of Ra-226 (3.7 x 1010 disintegrations per second, or 37 GBq).
Safety Handling of radium has been blamed for Marie Curie's premature death. • Radium is highly radioactive and its decay product, radon gas, is also radioactive. Since radium is chemically similar to calcium, it has the potential to cause great harm by replacing it in bones. Inhalation, injection, ingestion or body exposure to radium can cause cancer and other disorders. Stored radium should be ventilated to prevent accumulation of radon. • Emitted energy from the decay of radium ionizes gases, affects photographic plates, causes sores on the skin, and produces many other detrimental effects.
Radium
Further reading • Macklis, R. M. (1993). "The great radium scandal". Scientific American 269 (2): 94–99. • Clark, Claudia (1987). Radium Girls: Women and Industrial Health Reform, 1910–1935. University of North Carolina Press. ISBN ISBN 0-8078-4640-6.
See also • Decay chains • Radium Girls
References [1] " Mass Media & Environmental Conflict - Radium Girls (http:/ / www. radford. edu/ ~wkovarik/ envhist/ radium. html)". . Retrieved 2009-08-01. [2] " French Web site featuring products (medicines, mineral water, even underwear) containing radium (http:/ / www. dissident-media. org/ infonucleaire/ radieux. html)". . Retrieved 2009-08-01. [3] Cherbonnier, Alice (1997-10-01). " Nasal Radium Irradiation of Children Has Health Fallout (http:/ / baltimorechronicle. com/ rupnose. html)". Baltimore Chronicle. . Retrieved 2009-08-01. [4] Hayter, Charles (2005). " The Politics of Radon Therapy in the 1930s (http:/ / books. google. com/ books?id=NtKUdnjaCxMC& pg=PA135)". An Element of Hope: Radium and the Response to Cancer in Canada, 1900–1940. McGill-Queen's Press. ISBN 9780773528697. . [5] Pierre Curie, Madame Pierre Curie, and Gustave Bémont (1898). " Sur une nouvelle substance fortement radio-active, contenue dans la pechblende (On a new, strongly radioactive substance contained in pitchblende) (http:/ / www. aip. org/ history/ curie/ discover. htm)". Comptes Rendus 127: 1215–1217. . Retrieved 2009-08-01. [6] Marie Curie and André Debierne (1910). " Sur le radium métallique" (On metallic radium) (http:/ / visualiseur. bnf. fr/ CadresFenetre?O=NUMM-3104& I=523& M=tdm)" (in French). Comptes Rendus 151: 523–525. . Retrieved 2009-08-01. [7] J. J. Livingood (1936). "Deuteron-Induced Radioactivities". Phys Rev 50 (5): 425–434. doi: 10.1103/PhysRev.50.425 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 50. 425). [8] "Radium" (http:/ / periodic. lanl. gov/ elements/ 88. html), Los Alamos National Laboratory. Retrieved on 2009-08-05. [9] " Chart Nuclides by the National Nuclear Data Center (NNDC) (http:/ / www. nndc. bnl. gov/ chart/ reZoom. jsp?newZoom=3)". . Retrieved 2009-08-01.
• Albert Stwertka (1998). Guide to the Elements - Revised Edition. Oxford University Press. ISBN 0-19-508083-1. • Denise Grady (October 6, 1998). " A Glow in the Dark, and a Lesson in Scientific Peril (http:/ / www. nytimes. com/ library/ national/ science/ 100698sci-radium. html)". The New York Times. http:/ / www. nytimes. com/ library/ national/ science/ 100698sci-radium. html. Retrieved 2007-12-25. • Nanny Fröman (1 December 1996). " Marie and Pierre Curie and the Discovery of Polonium and Radium (http:/ / nobelprize. org/ nobel_prizes/ physics/ articles/ curie/ index. html)". Nobel Foundation. http:/ / nobelprize. org/ nobel_prizes/ physics/ articles/ curie/ index. html. Retrieved 2007-12-25.
6
Radium
External links • WebElements.com - Radium (http:/ / www. webelements. com/ webelements/ elements/ text/ Ra/ index. html) (also used as a reference) • Lateral Science - Radium Discovery (http:/ / www. lateralscience. co. uk/ radium/ RaDisc. html) • Photos of Radium Water Bath in Oklahoma (http:/ / www. markwshead. com/ stuffHappens/ radium. html) • NLM Hazardous Substances Databank – Radium, Radioactive (http:/ / toxnet. nlm. nih. gov/ cgi-bin/ sis/ search/ r?dbs+ hsdb:@term+ @na+ @rel+ radium,+ radioactive) • Reproduction of a 1942 comic book ad selling a "Radiumscope" to children (http:/ / www. lileks. com/ institute/ funny/ 07/ 40. html) pnb:میڈیر
7
Article Sources and Contributors
Article Sources and Contributors Radium Source: http://en.wikipedia.org/w/index.php?oldid=309100208 Contributors: 1toughnuke, 7, A Brave New World, A.Ou, Achaemenes, Adambro, Adashiel, Ahoerstemeier, Alansohn, AlimanRuna, Aloysius, Ampix0, Anaxial, Anclation, Andre Engels, Andrewa, Andy120290, Anonymous Dissident, Antandrus, Arkuat, Beetstra, Benbest, Bfigura's puppy, Big Brother 1984, Bkell, Bkonrad, Blainster, Bobo192, Bobwp, Brianhe, Bryan Derksen, Bubbachuck, CYD, Can't sleep, clown will eat me, CanadianLinuxUser, CanisRufus, Capecodeph, Capricorn42, Carsrac, Cbrown1023, Christoph Scholz, Clay66, Conversion script, Crazycomputers, Cryptic C62, Cureden, Cwkmail, Cyclonenim, DVD R W, Danny, Dannyc77, DariusMazeika, DarkAudit, Darrien, Darth Panda, Dave6, David Latapie, David spector, David.Monniaux, Dcb1995, Deadlord114, Dennis Brown, DennyColt, Dephillips21, DerHexer, Derek Ross, Discospinster, Djma12, Dr Zak, Dr.alf, DrBob, Dreamyshade, Drilnoth, Drini, Dryman, Dsfhsdkfjshdfkjlsdh, E9, Edgar181, El, El C, Emperorbma, Epbr123, Eric119, Erik Zachte, Evo584, Faiellie, Femto, Fibonacci, Flowerpotman, Fluri, Fonzy, Frankenpuppy, Gene Nygaard, Gilgamesh he, Gillyweed, Glenn, Gogo Dodo, Goldenchocolate, Greatpatton, Hak-kâ-ngìn, Hayabusa future, HazyM, Hdt83, Hqb, Humanist, IForgotToEatBreakFast, Icairns, Ideyal, Imjustmatthew, Iridescent, J.delanoy, JDspeeder1, JForget, JaGa, Jaraalbe, Javert, Jeronimo, JimVC3, Jmrwacko, Joanjoc, John, John C PI, Joshtynan, Joyous!, Julian Mendez, JunCTionS, Jurand, Kafka Liz, Kaiba, Kalamkaar, Karl-Henner, KeithD, Kesho, Kilo-Lima, King of Hearts, Kpalion, Kwamikagami, LarryMorseDCOhio, Leafyplant, Lightmouse, Likethesunshine, Lugnuts, Luk, MPerel, Malbi, Marc Venot, Martin451, Materialscientist, Mav, Mayooresan, McSly, Mdf, Mike2379, Mkweise, Moreschi, Muncadunc, Mwanner, Nakon, NawlinWiki, Nergaal, Neverquick, Nihiltres, NotALizard, Oatmeal batman, Obli, Onevalefan, Ossmann, Paul1953h, Philip Trueman, Polyparadigm, Poolkris, Poor Yorick, Postdlf, Primate, Pstudier, R, Ranveig, Razorflame, Rboatright, Remember, RexNL, Rich Farmbrough, Richard Arthur Norton (1958- ), Rifleman 82, Rjd0060, Robin Patterson, Rory096, Roudhound123, Rsm99833, Rursus, Sam8, Saperaud, Schneelocke, Science Focus, Scottmsg, Sengkang, Shawn in Montreal, Shinkolobwe, Shoeofdeath, Shotwell, Shuheziang, SimonP, Sjö, Sl, Smokefoot, Soliloquial, Soosed, SpookyMulder, Squiddy, Stack, Staffelde, Stifynsemons, Stone, Svante, SvenskaJohannes, Synchronism, T-Bone, Tadiew, Tagishsimon, TenOfAllTrades, Thebeast11, Thedjatclubrock, Thom.fynn, Tim Starling, Tobias Hoevekamp, Travis.Thurston, Travist, Triskaideka, Trumpet marietta 45750, Trusilver, Urbster1, Useight, Vargenau, Viriditas, Vsmith, Watch37264, Whosasking, Wikipeedio, Wikivendett, XJamRastafire, Xiahou, XoxoEmily, Xp54321, Yekrats, Yggdræsil, Yortzec, Yyy, Zaxgs, Zzorse, Zzyzx11, 491 anonymous edits
Image Sources, Licenses and Contributors image:Ra-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Ra-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Paddy, Saperaud
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
8
Actinium
1
Actinium radium ← actinium → thoriumLa ↑ Ac ↓ Ute
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HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
89Ac Periodic table
Appearance silvery General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacity Atomic properties Oxidation states ElectronegativityIonization energies 2nd: 1170 kJ·mol−1Covalent radius Miscellaneous Crystal structureMagnetic orderingThermal conductivityCAS registry number Most stable isotopes Main article: Isotopes of actinium iso
N.A.
half-life
DM
DE (MeV)
DP
225
syn
10 days
α
5.935
221
226
syn
29.37 hours
β−
1.117
226
ε
0.640
226
α
5.536
222
β−
0.045
227
α
5.042
223
Ac Ac
227
Ac
100%
21.773 years
Fr Th Ra Fr Th Fr
Actinium actinium, Ac, 89 actiniden/a, 7, f(227) g·mol−1 [Rn] 6d1 7s2 2, 8, 18, 32, 18, 9, 2 (Image) solid 10 g·cm−3 (circa) 1323 K,1050 °C,1922 °F 3471 K,3198 °C,5788 °F 14 kJ·mol−1 400 kJ·mol−1 (25 °C) 27.2 J·mol−1·K−1 3 (neutral oxide) 1.1 (Pauling scale) 1st: 499 kJ·mol−1215 pm face-centered cubic no data −1 −1 (300 K) 12 W·m ·K 7440-34-8 Actinium (pronounced /ækˈtɪniəm/) is a radioactive chemical element with the symbol Ac and atomic number 89, which was discovered in 1899. It was the first non-primordial radioactive element to be isolated, although polonium, radium and radon were observed before, but not isolated until 1902. It gave the name to the actinoid series, a group of 15 similar elements between actinium and lawrencium in the periodic table.
History Actinium was discovered in 1899 by André-Louis Debierne, a French chemist, who separated it from pitchblende as a substance being similar to titanium (1899)[1] or similar to thorium (1890).[2] Friedrich Oskar Giesel independently discovered actinium in 1902[3] as a substance being similar to lanthanum and called it "emanium" in 1904.[4] After a comparison of substances in 1904, Debierne's name was retained because it had seniority.[5] [6] The history of the discovery stayed questionable and in publications from 1971[7] and later in 2000[8] showed that the claims of André-Louis Debierne in 1904 conflict with the publications in 1899 and 1890. The word actinium comes from the Greek aktis, aktinos, meaning beam or ray.
Characteristics Actinium is a silvery, radioactive, metallic element. Due to its intense radioactivity, actinium glows in the dark with a pale blue light. The chemical behavior of actinium is similar to that of the rare earth element lanthanum.[9]
Chemistry Actinium shows similar chemical behavior to lanthanum. Due to this similarity the separation of actinium from lanthanum and the other rare earth elements, which are also present in uranium ores was difficult. Solvent extraction and ion exchange chromatography was used for the separation.[10] Only a limited amount of actinium compounds is known, for example AcF3, AcCl3, AcBr3, AcOF, AcOCl, AcOBr, Ac2S3, Ac2O3 and AcPO4. All the mentioned compounds are similar to the corresponding lanthanum compounds and shows that actinium compounds are generally in the oxidation state of +3.[11]
2
Actinium
Relationship with actinoids Actinium is the first element of the actinoids and gave the group its name, similar to lanthanum for the lanthanoids. The group of elements is more diverse than the lanthanoids and therefore it took until 1945 when Glenn T. Seaborg proposed the most significant change to Mendeleev's periodic table, by introducing the actinoids.
Isotopes Naturally occurring actinium is composed of 1 radioactive isotope; 227Ac. 36 radioisotopes have been characterized with the most stable being 227Ac with a half-life of 21.772 y, 225Ac with a half-life of 10.0 days, and 226Ac with a half-life of 29.37 h. All of the remaining radioactive isotopes have half-lives that are less than 10 hours and the majority of these have half-lives that are less than 1 minute. The shortest-lived isotope of actinium is 217Ac which decays through alpha decay and electron capture. It has a half-life of 69 ns. Actinium also has 2 meta states.[12] Purified 227Ac comes into equilibrium with its decay products at the end of 185 days, and then decays according to its 21.773-year half-life. The isotopes of actinium range in atomic weight from 206 u (206Ac) to 236 u (236Ac).[12]
Occurrence Actinium is found in trace amounts in uranium ore, but more commonly is made in milligram amounts by the neutron irradiation of 226Ra in a nuclear reactor. Actinium metal has been prepared by the reduction of actinium fluoride with lithium vapor at about 1100 to 1300°C.[9] Actinium is found only in traces in uranium ores as 227Ac, an α and β emitter with a half-life of 21.773 years. One ton of uranium ore contains about a tenth of a gram of actinium. The actinium isotope 227Ac is a transient member of the actinium series decay chain, which begins with the parent isotope 235U (or 239Pu) and ends with the stable lead isotope 207Pb. Another actinium isotope (225Ac) is transiently present in the neptunium series decay chain, beginning with 237Np (or 233U) and ending with near-stable bismuth (209Bi).
Applications It is about 150 times as radioactive as radium, making it valuable as a neutron source for energy. Otherwise it has no significant industrial applications.[13] 225
Ac is used in medicine to produce 213Bi in a reusable generator or can be used alone as an agent for radio-immunotherapy for Targeted Alpha Therapy (TAT).[14] 225Ac was first produced artificially by the Institute for Transuranium Elements (ITU) in Germany using a cyclotron and by Dr Graeme Melville at St George Hospital in Sydney using a linac in 2000.[15]
3
Actinium
4
Precautions 227
Ac is extremely radioactive, and in terms of its potential for radiation induced health effects[16] 227Ac is even more dangerous than plutonium. Ingesting even small amounts of 227 Ac would be fatal.
See also • Actinium series
External links • WebElements.com - Actinium [17] • NLM Hazardous Substances Databank – Actinium, Radioactive
[18]
pnb:مینیٹکیا
References [1] Debierne, André-Louis (1899). " Sur un nouvelle matière radio-active (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k3085b/ f593. table)". Comptes rendus 129: 593–595. . [2] Debierne, André-Louis (1900-1901). " Sur un nouvelle matière radio-actifl'actinium (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k3086n/ f906. table)". Comptes rendus 130: 906–908. . [3] Giesel, Friedrich Oskar (1902). "Ueber Radium und radioactive Stoffe". Berichte der Deutschen Chemische Geselschaft 35 (3): 3608–3611. doi: 10.1002/cber.190203503187 (http:/ / dx. doi. org/ 10. 1002/ cber. 190203503187). [4] Giesel, Friedrich Oskar (1904). "Ueber den Emanationskörper (Emanium)". Berichte der Deutschen Chemische Geselschaft 37 (2): 1696–1699. doi: 10.1002/cber.19040370280 (http:/ / dx. doi. org/ 10. 1002/ cber. 19040370280). [5] Giesel, Friedrich Oskar (1904). "Ueber Emanium". Berichte der Deutschen Chemische Geselschaft 37 (2): 1696–1699. doi: 10.1002/cber.19040370280 (http:/ / dx. doi. org/ 10. 1002/ cber. 19040370280). [6] Giesel, Friedrich Oskar (1905). "Ueber Emanium". Berichte der Deutschen Chemische Geselschaft 38 (1): 775–778. doi: 10.1002/cber.190503801130 (http:/ / dx. doi. org/ 10. 1002/ cber. 190503801130). [7] Kirby, H. W. (1971). " The Discovery of Actinium (http:/ / www. jstor. org/ stable/ view/ 229943?seq=1)". Isis 62 (3): 290–308. doi: 10.1086/350760 (http:/ / dx. doi. org/ 10. 1086/ 350760). . [8] Adloff, J. P. (2000). "The centenary of a controversial discovery: actinium". Radiochim. Acta, 88: 123–128. doi: 10.1524/ract.2000.88.3-4.123 (http:/ / dx. doi. org/ 10. 1524/ ract. 2000. 88. 3-4. 123). [9] Stites, Joseph G.; Salutsky, Murrell L. Stone, Bob D. (1955). "Preparation of Actinium Metal". J. Am. Chem. Soc. 77 (1): 237–240. doi: 10.1021/ja01606a085 (http:/ / dx. doi. org/ 10. 1021/ ja01606a085). [10] Katz, J. J.; Manning, W. M. (1952). "Chemistry of the Actinide Elements Annual Review of Nuclear Science". Annual Review of Nuclear Science 1: 245–262. doi: 10.1146/annurev.ns.01.120152.001333 (http:/ / dx. doi. org/ 10. 1146/ annurev. ns. 01. 120152. 001333). [11] Sherman, Fried; Hagemann, French; Zachariasen, W. H. (1950). "The Preparation and Identification of Some Pure Actinium Compounds". Journal of the American Chemical Society 72: 771–775. doi: 10.1021/ja01158a034 (http:/ / dx. doi. org/ 10. 1021/ ja01158a034). [12] Audi, Georges (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3–128. doi: 10.1016/j.nuclphysa.2003.11.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 001). [13] Dixon, W.R.; Bielesch, Alice; Geiger K.W. (1957). " Neutron Spectrum of an Actinium–Beryllium Source (http:/ / pubs. nrc-cnrc. gc. ca/ cgi-bin/ rp/ rp2_abst_e?cjp_p57-075_35_ns_nf_cjp)". Can. J. Phys./Rev. Can. Phys. 35 (6): 699–702. . [14] Bolla, Rose A.; Malkemusa, Dairin; Mirzadeh, Saed (2005). "Production of actinium-225 for alpha particle mediated radioimmunotherapy". Applied Radiation and Isotopes 62 (5): 667–679. doi: 10.1016/j.apradiso.2004.12.003 (http:/ / dx. doi. org/ 10. 1016/ j. apradiso. 2004. 12. 003). [15] Melville, G; Allen, Bj (Apr 2009). "Cyclotron and linac production of Ac-225.". Applied radiation and isotopes : including data, instrumentation and methods for use in agriculture, industry and medicine 67 (4): 549–55. doi: 10.1016/j.apradiso.2008.11.012 (http:/ / dx. doi. org/ 10. 1016/ j. apradiso. 2008. 11. 012). ISSN 0969-8043 (http:/ / worldcat. org/ issn/ 0969-8043). PMID 19135381.
Actinium [16] Langham, W.; Storer, J. (1952). "Toxicology of Actinium Equilibrium Mixture". Los Alamos Scientific Lab.: Technical Report. doi: 10.2172/4406766 (http:/ / dx. doi. org/ 10. 2172/ 4406766). [17] http:/ / www. webelements. com/ webelements/ elements/ text/ Ac/ index. html [18] http:/ / toxnet. nlm. nih. gov/ cgi-bin/ sis/ search/ r?dbs+ hsdb:@term+ @na+ @rel+ actinium,+ radioactive
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Article Sources and Contributors
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6
Thorium
1
Thorium actinium ← thorium → protactinium Ce ↑ Th ↓ (Uqn) Periodic Table Extended Periodic Table General Name, symbol, number
thorium, Th, 90
Element category
Actinides
Group, Period, Block
n/a, 7, f
Appearance
silvery white Standard atomic weight
−1
232.0381(2) g·mol
Electron configuration
2
2
[Rn] 6d 7s
Electrons per shell
2, 8, 18, 32, 18, 10, 2 Physical properties
Phase
solid Density (near r.t.)
11.7 g·cm
Melting point
−3
2115 K (1842 °C, 3348 °F) Boiling point
5061 K (4788 °C, 8650 °F)
Heat of fusion
13.81 kJ·mol−1
Heat of vaporization
514 kJ·mol−1
Specific heat capacity
(25 °C) 26.230 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
2633
2907
3248
3683
4259
5055
Atomic properties Crystal structure Oxidation states Electronegativity
face centered cubic 4, 3, 2 (weakly basic oxide) 1.3 (Pauling scale)
Thorium
2
Ionization energies (more)
1st: 587 kJ·mol−1 2nd: 1110 kJ·mol−1 3rd: 1930 kJ·mol−1
Atomic radius
179 pm
Covalent radius
206±6 pm Miscellaneous [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(0 °C) 147 nΩ·m
Thermal conductivity
(300 K) 54.0 W·m−1·K−1
Thermal expansion
(25 °C) 11.0 µm·m
Speed of sound (thin rod)
(20 °C) 2490 m/s
Young's modulus
79 GPa
Shear modulus
31 GPa
−1
Bulk modulus
−1
·K
54 GPa
Poisson ratio
0.27
Mohs hardness
3.0
Vickers hardness
350 MPa
Brinell hardness
400 MPa
CAS registry number
7440-29-1 Most-stable isotopes
Main article: Isotopes of thorium iso
NA
half-life
DM
DE (MeV)
DP
228
syn
1.9116 years
α
5.520
224
229
syn
7340 years
α
5.168
225
230
syn
75380 years
α
4.770
226
231
trace
25.5 hours
β
0.39
231
1.405×1010 years
α
4.083
228
24.1 days
β
0.27
234
Th Th Th Th
232
Th
234
Th
100% trace
Ra Ra Ra Pa Ra Pa
References
Thorium (pronounced /ˈθɔəriəm/) is a chemical element with the symbol Th and atomic number 90. It is a naturally occurring, slightly radioactive metal, which has been successfully used as an alternative nuclear fuel to uranium in the molten-salt reactor experiment (MSR) for several years to produce thermal energy. Thorium is abundant on Earth and this type of reactor can be built to operate significantly cleaner than uranium based power plants as the waste products are much easier to handle.[2] The benefits of the MSR are similar (fuel abundancy and clean operation) to what nuclear fusion promises.[3]
Thorium
Characteristics Physical Pure thorium is a silvery-white metal which is air-stable and retains its luster for several months. When contaminated with the oxide, thorium slowly tarnishes in air, becoming gray and finally black. The physical properties of thorium are greatly influenced by the degree of contamination with the oxide. The purest specimens often contain several tenths of a percent of the oxide. Pure thorium is soft, very ductile, and can be cold-rolled, swaged, and drawn. Thorium is dimorphic, changing at 1400 °C from a face-centered cubic to a body-centered cubic structure. Powdered thorium metal is often pyrophoric and should be carefully handled. When heated in air, thorium metal turnings ignite and burn brilliantly with a white light. Thorium has the largest liquid range of any element: 2946 °C between the melting point and boiling point. [4]
Chemical Thorium is slowly attacked by water, but does not dissolve readily in most common acids, except hydrochloric.[4] It dissolves in concentrated nitric acid containing a small amount of catalytic fluoride ion.[5]
Compounds Thorium compounds are stable in the +4 oxidation state.[6] Thorium dioxide has the highest melting point (3300 °C) of all oxides.[7] Thorium(IV) nitrate and thorium(IV) fluoride are known in their hydrated forms: Th(NO3)4·4H2O and ThF4·4H2O, respectively. The thorium center has square planar geometry.[6] Thorium(IV) carbonate, Th(CO3)2, is also known.[6] When treated with potassium fluoride and hydrofluoric acid, Th4+ forms the complex anion ThF62−, which precipitates as an insoluble salt, K2ThF6.[5] Thorium(IV) hydroxide, Th(OH)4, is highly insoluble in water, and is not amphoteric. The peroxide of thorium is rare in being an insoluble solid. This property can be utilized to separate thorium from other ions in solution.[5] In the presence of phosphate anions, Th4+ forms precipitates of various compositions, which are insoluble in water and acid solutions.[5]
Isotopes Naturally occurring thorium is composed mainly of one isotope: 232Th. 230Th occurs as the daughter product of 238U decay. Twenty-seven radioisotopes have been characterized, with the most abundant and/or stable being 232Th with a half-life of 14.05 billion years, 230Th with a half-life of 75,380 years, 229Th with a half-life of 7340 years, and 228Th with a half-life of 1.92 years. All of the remaining radioactive isotopes have half-lives that are less than thirty days and the majority of these have half-lives that are less than ten minutes. One isotope, 229Th, has a nuclear isomer (or metastable state) with a remarkably low excitation energy of 7.6 eV.[8] The known isotopes of thorium range in atomic weight from 210 u (210Th) to 236 u (236Th).[9]
3
Thorium
Applications Applications of thorium:[4] • Thorium is used as an alloying element in magnesium, used in aircraft engines, imparting high strength and creep resistance at elevated temperatures.[10] • Thorium is also used as an alloying agent in gas tungsten arc welding (GTAW) to increase the melting temperature of tungsten electrodes and improve arc stability. The electrodes labeled EWTH-1 contain 1% thorium, while the EWTH-2 contain 2%.[11] • Thorium is used to coat tungsten wire used in electronic equipment, improving the electron emission of heated cathodes. • Uranium-thorium age dating has been used to date hominid fossils. • Thorium is used as a fertile material for producing nuclear fuel. In particular, the proposed energy amplifier reactor design would employ thorium. Since thorium is more abundant than uranium, some nuclear reactor designs incorporate thorium in their fuel cycle. • Thorium may also be used directly as nuclear fuel instead of uranium, producing less transuranic waste. • Thorium is a very effective radiation shield, although it has not been used for this purpose as much as lead or depleted uranium. Applications of thorium dioxide (ThO2): • Mantles in portable gas lights. These mantles glow with a dazzling light (unrelated to radioactivity) when heated in a gas flame. • Used to control the grain size of tungsten used for electric lamps. • Used in heat-resistant ceramics like high-temperature laboratory crucibles. • Added to glass, it helps create glasses of a high refractive index and with low dispersion. Consequently, they find application in high-quality lenses for cameras and scientific instruments. • Has been used as a catalyst: • In the conversion of ammonia to nitric acid. • In petroleum cracking. • In producing sulfuric acid. • Thorium dioxide is the active ingredient of Thorotrast, which was used as part of X-ray diagnostics. This use has been abandoned due to the carcinogenic nature of Thorotrast.
Thorium as a nuclear fuel Thorium, as well as uranium and plutonium, can be used as fuel in a nuclear reactor. Although not fissile itself, 232Th will absorb slow neutrons to produce 233U, which is fissile. Hence, like 238U, it is fertile. Theoretically thorium is a more suitable fuel source than uranium. It is at least 4-5 times more abundant in nature than all of uranium isotopes combined and is fairly evenly spread around Earth, with many countries having large supplies of it. Also, preparation of thorium fuel does not require difficult and expensive enrichment process. The thorium fuel cycle creates mainly Uranium-233 which can be used for making nuclear weapons, and since there are no neutrons from spontaneous fission of U-233, U-233 can be used easily in a gun-type nuclear bomb[12] Thorium can and has been used to power nuclear energy plants using both the modified traditional Generation III reactor design and prototype Generation IV reactor designs.
4
Thorium When using thorium in modified light water reactor (LWR) problems include: the undeveloped technology for fuel fabrication; in traditional, once-through LWR designs potential problems in recycling thorium due to highly radioactive 228Th; some weapons proliferation risk due to production of 233U; and the technical problems (not yet satisfactorily solved) in reprocessing. Much development work is still required before the thorium fuel cycle can be commercialized for use in LWR, and the effort required seems unlikely while (or where) abundant uranium is available. Nevertheless, the thorium fuel cycle, with its potential for breeding fuel without fast neutron reactors, holds considerable potential long-term benefits. Thorium is significantly more abundant than uranium, and is a key factor in sustainable nuclear energy. Perhaps more importantly, thorium produces several orders of magnitude less long-lived radioactive waste. One of the earliest efforts to use a thorium fuel cycle took place at Oak Ridge National Laboratory in the 1960s. An experimental reactor was built based on MSR technology to study the feasibility of such an approach, using thorium-fluoride salt kept hot enough to be liquid, thus eliminating the need for fabricating fuel elements. This effort culminated in the Molten-Salt Reactor Experiment that used 232Th as the fertile material and 233U as the fissile fuel. This reactor has been operated successfully for about five years. However due to a lack of funding, the MSR program was discontinued in 1976. Nowadays this design is considered as Generation IV reactor. India's Kakrapar-1 reactor is the world's first reactor which utilizes thorium rather than depleted uranium to achieve power flattening across the reactor core.[13] India, which has about 25% of the world's thorium reserves, is developing a 300 MW prototype of a thorium-based Advanced Heavy Water Reactor (AHWR). The prototype is expected to be fully operational by 2011, following which five more reactors will be constructed.[14] India currently envisages to meet 30% of its electricity demand through thorium-based reactors by 2030.[15] In 2007, Norway was debating whether or not to focus on thorium plants, due to the existence of large deposits of thorium ores in the country, particularly at Fensfeltet, near Ulefoss in Telemark county. The primary fuel of the HT3R Project near Odessa, Texas, USA will be ceramic-coated thorium beads.
History M. T. Esmark found a black mineral on Løvøy Island, Norway and gave a sample to Professor Jens Esmark, a noted mineralogist who was not able to identify it, so he sent a sample to the Swedish chemist Jöns Jakob Berzelius for examination in 1828.[16] [17] [18] Berzelius analyzed it and named it after Thor, the Norse god of thunder. The metal had virtually no uses until the invention of the gas mantle in 1885. In 1898 thorium was first observed to be radioactive, independently, by Polish-French physicist Marie Curie and English chemist Gerhard Carl Schmidt.[19] [20] [21] Between 1900 and 1903, Ernest Rutherford and Frederick Soddy showed how thorium decayed at a fixed rate over time into a series of other elements. This observation led to the identification of half life as one of the outcomes of the alpha particle experiments that led to their disintegration theory of radioactivity.[22]
5
Thorium
6
The crystal bar process (or Iodide process) was discovered by Anton Eduard van Arkel and Jan Hendrik de Boer in 1925 to produce high-purity metallic thorium.[23] The name ionium was given early in the study of radioactive elements to the 230Th isotope produced in the decay chain of 238U before it was realized that ionium and thorium were chemically identical. The symbol Io was used for this supposed element.
Occurrence
Monazite, a rare-earth-and-thorium phosphate mineral, is the primary source of the world's thorium
Thorium is found in small amounts in most rocks and soils, where it is about four times more abundant than uranium, and is about as common as lead. Soil commonly contains an average of around 12 parts per million (ppm) of thorium. Thorium occurs in several minerals including thorite (ThSiO4), thorianite (ThO2 + UO2) and monazite. The latter is most common and may contain up to about 12% thorium oxide. Thorium-containing monazite(Ce) occurs in Africa, Antarctica, Australia, Europe, India, North America, and South America.[4] [24] 232
Th decays very slowly (its half-life is comparable to the age of the Universe) but other thorium isotopes occur in the thorium and uranium decay chains. Most of these are short-lived and hence much more radioactive than 232Th, though on a mass basis they are negligible.
Thorium extraction
Thorium has been extracted chiefly from monazite through a complex multi-stage process. The monazite sand is dissolved in hot concentrated sulfuric acid (H2SO4). Thorium is extracted as an insoluble residue into an organic phase containing an amine. Next it is separated or "stripped" using an ion such as nitrate, chloride, hydroxide, or carbonate, returning the thorium to an aqueous phase. Finally, the thorium is precipitated and collected.[25] Several methods are available for producing thorium metal: it can be obtained by reducing thorium oxide with calcium, by electrolysis of anhydrous thorium chloride in a fused mixture of sodium and potassium chlorides, by calcium reduction of thorium tetrachloride mixed with anhydrous zinc chloride, and by reduction of thorium tetrachloride with an
Thorium
7
alkali metal.[4]
Distribution Present knowledge of the distribution of thorium resources is poor because of the relatively low-key exploration efforts arising out of insignificant demand.[26] There are two sets of estimates that define world thorium reserves, one set by the US Geological Survey (USGS) and the other supported by reports from the OECD and the International Atomic Energy Agency (the IAEA). Under the USGS estimate, Australia and India have particularly large reserves of thorium. India and Australia are believed to possess about 300,000 metric tonnes each; i.e. each country possessing 25% of the world's thorium reserves.[27] However, in the OECD reports, estimates of Australian's Reasonably Assured Reserves (RAR) of Thorium indicate only 19,000 metric tonnes and not 300,000 tonnes as indicated by USGS. The two sources vary wildly for countries such as Brazil, Turkey, and Australia. However, both reports appear to show some consistency with respect to India's thorium reserve figures, with 290,000 metric tonnes (USGS) and 319,000 metric tonnes (OECD/IAEA). Furthermore the IAEA report mentions that India possesses two thirds (67%) of global reserves of monazite, the primary thorium ore. The IAEA also states that recent reports have upgraded India's thorium deposits up from approximately 300,000 metric tonnes to 650,000 metric tonnes.[28] Therefore, the IAEA and OECD appear to conclude that Brazil and India may actually possess the lion's share of world's thorium deposits. • The prevailing estimate of the economically available thorium reserves comes from the US Geological Survey, Mineral Commodity Summaries (1997-2006):[29] [30] Country
Th Reserves (tonnes)
Th Reserve Base (tonnes)
Australia
300,000
340,000
India
290,000
300,000
Norway
170,000
180,000
United States
160,000
300,000
Canada
100,000
100,000
South Africa
35,000
39,000
Brazil
16,000
18,000
Malaysia
4,500
4,500
Other Countries
95,000
100,000
World Total
1,200,000
1,400,000
Note: The Australian figures are based on assumptions and not on actual geological surveys, therefore the figures cited for Australia may be misleading, should be treated with caution and could possibly indicate inflated values for Australia's actual reserves of thorium; note the OECD estimates of Australian's Reasonably Assured Reserves (RAR) of Thorium (listed below) indicate only 19,000 metric tonnes and not 300,000 tonnes as listed above. • Another estimate of Reasonably Assured Reserves (RAR) and Estimated Additional Reserves (EAR) of thorium comes from OECD/NEA, Nuclear Energy, "Trends in Nuclear Fuel Cycle", Paris, France (2001):[31] Country
RAR Th (tonnes)
EAR Th (tonnes)
Thorium
8
Brazil
606,000
700,000
Turkey
380,000
500,000
India
319,000
—
United States
137,000
295,000
Norway
132,000
132,000
Greenland
54,000
32,000
Canada
45,000
128,000
Australia
19,000
—
South Africa
18,000
—
Egypt
15,000
309,000
Other Countries
505,000
—
World Total
2,230,000
2,130,000
Precautions See Actinides in the environment for details of the environmental aspects of thorium. Powdered thorium metal will often ignite spontaneously in air (it is pyrophoric) and should be handled carefully. Natural thorium decays very slowly compared to many other radioactive materials, and the alpha radiation emitted cannot penetrate human skin. Owning and handling small amounts of thorium, such as a gas mantle, is considered safe if care is taken not to ingest the thorium—lungs and other internal organs can be penetrated by alpha radiation. Exposure to an aerosol of thorium can lead to increased risk of cancers of the lung, pancreas and blood. Exposure to thorium internally leads to increased risk of liver diseases. This element has no known biological role. See also Thorotrast.
See also • Periodic table • Nuclear reactor • Decay chain • Sylvania Electric Products explosion
External links WebElements.com — Thorium [32] The World Nuclear Association [33] European Nuclear Society — Natural Decay Chains [34] often-quoted article by Michael Anissimov advocating adopting Thorium reactors [35] Thorium information page [36] New Age Nuclear: article on thorium reactors | Cosmos Magazine [37] ATSDR ToxFAQs — Thorium [38] Thorium as a Secure Nuclear Fuel Alternative [39] The Endless Refrigerator/Freezer Deodorizer [40], a commercial product which claimed to destroy odours 'forever.' Made with thorium-232. • Is thorium the answer to our energy crisis? [41] • Thorium Energy [42] Blog, discussion forum and document repository • • • • • • • • •
Thorium • Another thorium information page
9 [43]
pnb:میروھت
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] " Thorium (http:/ / www. world-nuclear. org/ info/ inf62. html)". World Nuclear Association. . Retrieved 2009-08-08. [3] " Thorium Power Technology (http:/ / www. thoriumpower. com/ default2. asp?nav=technology_solutions)". . Retrieved 2009-08-08. [4] C. R. Hammond (2004). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304857. [5] Earl K. Hyde (1960). The radiochemistry of thorium (http:/ / www. radiochemistry. org/ periodictable/ pdf_books/ pdf/ rc000034. pdf). Subcommittee on Radiochemistry, National Academy of Sciences—National Research Council. . [6] " Toxicological Profile Information Sheet (http:/ / www. atsdr. cdc. gov/ toxprofiles/ tp147-c3. pdf)". Department of Health and Human Services. . Retrieved 2009-05-21. [7] Emsley, John (2001). Nature's Building Blocks ((Hardcover, First Edition) ed.). Oxford University Press. pp. 441. ISBN 0198503407. [8] B. R. Beck et al. (2007). "Energy Splitting of the Ground-State Doublet in the Nucleus 229Th". Phys. Rev. Lett. 98: 142501. doi: 10.1103/PhysRevLett.98.142501 (http:/ / dx. doi. org/ 10. 1103/ PhysRevLett. 98. 142501). [9] J. Uusitalo et al. (1995). "α decay of the new isotopes 210Th and 211Th". Phys. Rev. C 52: 113. doi: 10.1103/PhysRevC.52.113 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 52. 113). [10] ed. by Michael M. Avedesian, Prepared under the direction of the ASM International Handbook Committee. (1999). " Microstructure of Magnesium and Magnesium Alloys (http:/ / books. google. de/ books?id=0wFMfJg57YMC& pg=PA28)". Magnesium and magnesium alloys. Materials Park, OH: ASM International. p. 28. ISBN 9780871706577. . [11] Larry Jeffus. (2003). " Types of Tungsten (http:/ / books. google. de/ books?id=zeRiW7en7HAC& pg=RA1-PA750)". Welding : principles and applications. Clifton Park, N.Y.: Thomson/Delmar Learning. p. 350. ISBN 9781401810467. . [12] R. Wilson (1998). " Accelerator Driven Subcritical Assemblies (http:/ / phys4. harvard. edu/ ~wilson/ publications/ ppaper703. html)". Report to Energy Environment and Economy Committee, U.S. Global Strategy Council. . [13] " Thorium: Cleaner Nuclear Power? (http:/ / www. power-technology. com/ features/ feature1141/ )". . [14] " Development work on 300 MW advanced heavy water reactor at advanced stage (http:/ / timesofindia. indiatimes. com/ articleshow/ 3864684. cms)". . [15] " Indian Thorium based reactor design complete (http:/ / www. indiadaily. com/ editorial/ 19093. asp)". . [16] " Thorium (http:/ / www. bbc. co. uk/ dna/ h2g2/ A3768861)". BBC.co. . Retrieved 2007-01-18. [17] J. J. Berzelius (1829). " Untersuchung eines neues Minerals und einer darin erhalten zuvor unbekannten Erde (Investigation of a new mineral and of a previously unknown earth contained therein) (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k151010. pleinepage. r=Annalen+ der+ Physic. f395. langFR)". Annalen der Physik und Chemie 16: 385-415. . (modern citation: Annalen der Physik, vol. 92, no. 7, pages 385-415) [18] J. J. Berzelius (1829). "Undersökning af ett nytt mineral (Thorit), som innehåller en förut obekant jord" (Investigation of a new mineral (thorite), as contained in a previously unknown earth)". Kungliga Svenska Vetenskaps Akademiens Handlingar (Transactions of the Royal Swedish Science Academy): 1-30. [19] Marie Curie (1898). "Rayons émis par les composés de l'uranium et du thorium (Rays emitted by compounds of uranium and thorium)". Comptes Rendus 126: 1101-1103. [20] G. C. Schmidt (1898). "Über die vom Thorium und den Thoriumverbindungen ausgehende Strahlung (On the radiation emitted by thorium and thorium compounds)". Verhandlungen der Physikalischen Gesellschaft zu Berlin (Proceedings of the Physical Society in Berlin) 17: 14-16. [21] G. C. Schmidt (1898). " Über die von den Thorverbindungen und einigen anderen Substanzen ausgehende Strahlung (On the radiation emitted by thorium compounds and some other substances) (http:/ / gallica. bnf. fr/ ark:/ 12148/ bpt6k153068. image. r=Annalen+ der+ Physic. f149. langFR)". Annalen der Physik und Chemie 65: 141-151. . (modern citation: Annalen der Physik, vol. 301, pages 141-151 (1898)). [22] Simmons, John Galbraith (1996). The Scientific 100. Seacaucus NJ: Carol. p. 19. [23] van Arkel, A.E.; de Boer, J.H. (1925). "Preparation of pure titanium, zirconium, hafnium, and thorium metal". Zeitschrift für Anorganische und Allgemeine Chemie 148: 345–350.
Thorium [24] " Monazite-(Ce): Monazite-(Ce) mineral information and data (http:/ / www. mindat. org/ min-2751. html)". . Retrieved 18 May 2009. [25] Crouse, David (1959). "The Amex Process for Extracting Thorium Ores with Alkyl Amines". Industrial & Engineering Chemistry 51: 1461. doi: 10.1021/ie50600a030 (http:/ / dx. doi. org/ 10. 1021/ ie50600a030). [26] K.M.V. Jayaram. " An Overview of World Thorium Resources, Incentives for Further Exploration and Forecast for Thorium Requirements in the Near Future (http:/ / www. iaea. org/ inis/ aws/ fnss/ fulltext/ 0412_1. pdf)". . [27] " US approves Indian nuclear deal (http:/ / news. bbc. co. uk/ 2/ hi/ south_asia/ 6219998. stm)". BBC News. 2006-12-09. . [28] IAEA: Thorium fuel cycle — Potential benefits and challenges (http:/ / www-pub. iaea. org/ MTCD/ publications/ PDF/ TE_1450_web. pdf). pp. 45. . [29] " U.S. Geological Survey, Mineral Commodity Summaries - Thorium (http:/ / minerals. usgs. gov/ minerals/ pubs/ commodity/ thorium/ index. html#mcs)". . [30] " Information and Issue Briefs - Thorium (http:/ / www. world-nuclear. org/ info/ inf62. htm)". World Nuclear Association. . Retrieved 2006-11-01. [31] IAEA: Thorium fuel cycle — Potential benefits and challenges (http:/ / www-pub. iaea. org/ MTCD/ publications/ PDF/ TE_1450_web. pdf). pp. 45(table 8), 97(ref 78). . [32] [33] [34] [35] [36] [37]
http:/ / www. webelements. com/ webelements/ elements/ text/ Th/ index. html http:/ / www. world-nuclear. org/ http:/ / www. euronuclear. org/ info/ encyclopedia/ d/ decaybasinnatural. htm http:/ / www. acceleratingfuture. com/ michael/ blog/ 2006/ 10/ a-nuclear-reactor-in-every-home/ http:/ / www. world-nuclear. org/ info/ inf62. htm http:/ / www. cosmosmagazine. com/ node/ 348/
[38] http:/ / www. atsdr. cdc. gov/ tfacts147. html [39] http:/ / www. ensec. org/ index. php?option=com_content& view=article& id=187:thorium-as-a-secure-nuclear-fuel-alternative& catid=94:0409content& Itemid=342 [40] [41] [42] [43]
http:/ / www. orau. org/ ptp/ collection/ quackcures/ endless. htm http:/ / news. independent. co. uk/ sci_tech/ article2070374. ece http:/ / thoriumenergy. blogspot. com http:/ / www. energyfromthorium. com/
10
Article Sources and Contributors
Article Sources and Contributors Thorium Source: http://en.wikipedia.org/w/index.php?oldid=309574608 Contributors: 2mcm, 65.68.87.xxx, Aarchiba, Acalamari, Acid88, Adapter, Ahoerstemeier, Alansohn, AlexW, Alterrabe, AndonicO, Andres, Andrewa, Antandrus, Archimerged, Arg, Arkuat, Audunv, Awatral, Bank top, Barneyg, Bayou Banjo, Beetstra, BenAlbahari, Benbest, BlueEarth, Bobo192, Borislav Dopudja, Brian Huffman, Brisvegas, Bryan Derksen, CAPS LOCK, CJGB, Cadmium, Cameronw22, Canthusus, Carnildo, Chill doubt, Chris 73, Chris Roy, Chriscf, Christine T, Chuckiesdad, Cmacd123, CommonsDelinker, Conversion script, Cryptic, Crzrussian, Cwkmail, Cyberevil, DMahalko, Danpat, Darrien, David Latapie, David R. Ingham, DavidMIA, Deor, Dewan357, Donarreiskoffer, DrHonzik, Dwmyers, Edgar181, El C, Emperorbma, Femto, Finefellow, Fionaclee, Foo1942, Fresheneesz, Fundamental metric tensor, Furrykef, G-Man, Gaius Cornelius, GeorgeTSLC, Graeme Bartlett, Graibeard, Greatpatton, Grey Shadow, Gsmcolect, Gthb, Guthrie, Hak-kâ-ngìn, Hallpriest9, Helge Skjeveland, Hersfold, Hinakana, Hitssquad, Holy Ganga, Icairns, Ideyal, Idleguy, Itub, J.delanoy, Janke, Jaraalbe, Jayshuler, JdH, Jdurg, Jim62sch, Joanjoc, John, Jons63, Julian Mendez, JunCTionS, KFSorensen, Kaihsu, Karelj, Kayleigheliz, Keenan Pepper, Keine entschuldigung, Kelovy, King Zebu, Kjramesh, Ktsquare, Kudz75, Kurykh, Kwamikagami, Kwilbur, Lamlott, LarryMorseDCOhio, LorenzoB, Malcolm Farmer, Marc Venot, Mark.murphy, Materialscientist, Mattisse, Mav, Megan1967, Michelet, Mithridates, Mnyaseen, Moogsi, Mr0t1633, Mrintel, Nailedtooth, Nakkiel, Neil916, Nergaal, Neverquick, Nick Y., Nihiltres, Ohnoitsjamie, Pakaran, Paradigm, PeepP, Peter Greenwell, Peter.thejackos, Peter8212991, PeterJeremy, PhasmatisNox, PiersOffshore, Piperh, Plunkey, Polonium, Poolkris, Pstudier, REQC, RamanVirk, Rbakker99, Rcnet, Remember, RexNL, Rgoodermote, Rich Farmbrough, Richard W.M. Jones, Rickterp, Riick, Rmhermen, Robert Merkel, Romanm, Royk, Ryan Lonswell, SDC, Santăr, Saperaud, Schneelocke, Scott Ritchie, Sfuerst, Sgokoluk, Shaddack, Shadowjams, Sholt, Shyam, Sillybilly, Sinneed, Sl, Sleigh, Slidersv, Soliloquial, Spangineer, Srikeit, Sstidman, Stifynsemons, Stone, Svante, Tagishsimon, Tempodivalse, Tetracube, The Haunted Angel, TheDestitutionOfOrganizedReligion, TheoClarke, Thingg, Thoreaulylazy, Thorml, Tobias Hoevekamp, Trevorloflin, Triforce of Power, Ummit, Versus22, Volland, Vsmith, Wachholder0, Walkerma, Warut, Whkoh, Whosasking, Wikieditor06, Wwoods, Yekrats, Yyy, Zerpent, Zoomzoom316, 270 anonymous edits
Image Sources, Licenses and Contributors image:Th-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Th-TableImage.png License: GNU Free Documentation License Contributors: Paddy, Paginazero, Saperaud Image:MonaziteUSGOV.jpg Source: http://en.wikipedia.org/w/index.php?title=File:MonaziteUSGOV.jpg License: Public Domain Contributors: Saperaud Image:Monazit opening acid.gif Source: http://en.wikipedia.org/w/index.php?title=File:Monazit_opening_acid.gif License: unknown Contributors: User:Hermann Luyken
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11
Protactinium
1
Protactinium thorium ← protactinium → uranium Pr ↑ Pa ↓ (Uqu) Periodic Table Extended Periodic Table General Name, symbol, number Element category Group, Period, Block Appearance
protactinium, Pa, 91 actinides n/a, 7, f bright, silvery metallic luster
Standard atomic weight Electron configuration Electrons per shell
−1
231.03588(2) g·mol 2
1
2
[Rn] 7s 6d 5f
2, 8, 18, 32, 20, 9, 2 Physical properties
Phase
solid Density (near r.t.)
Melting point
−3
15.37 g·cm
1841 K (1568 °C, 2854 °F) Boiling point
? 4300 K (? 4027 °C, ? 7280 °F)
Heat of fusion
12.34 kJ·mol−1
Heat of vaporization
481 kJ·mol−1 Atomic properties
Crystal structure Oxidation states
orthorhombic 2, 3, 4, 5 (weakly basic oxide)
Electronegativity Ionization energies
1.5 (Pauling scale) 1st: 568 kJ/mol
Atomic radius
163 pm
Covalent radius
200 pm Miscellaneous
Magnetic ordering
[1]
paramagnetic
Protactinium
2
Electrical resistivity
(0 °C) 177 nΩ·m
Thermal conductivity
(300 K) 47 W·m−1·K−1
CAS registry number
7440-13-3 Most-stable isotopes
Main article: Isotopes of protactinium iso
NA
half-life
DM
DE (MeV)
DP
229
syn
1.4 d
α
5.58
225
230
syn
17.4 d
ε
1.310
230
β−
0.563
230
32760 y
α
5.149
227
1.31 d
β
−
0.31
232
−
0.571
233
−
Pa Pa
231
Pa
~100%
Ac Th U Ac
232
syn
233
syn
26.967 d
β
syn
1.17 min
β
2.29
234
IT
0.0694
234
0.23
234
Pa Pa
234m
Pa
234
Pa
syn
6.75 h
−
β
U U U → Pa
U
References
Protactinium (pronounced /ˌproʊtækˈtɪniəm/) is a chemical element with the symbol Pa and atomic number 91. Its longest-lived and only naturally-occurring isotope, Pa-231, is a decay product of uranium-235 (U-235), and it has a half-life of 32,760 years.
Characteristics Protactinium is a metallic element that belongs to the actinoid group, with a bright metallic luster that it retains for some time in contact with air.[2] Protactinium is superconductive at temperatures below 1.4 K.[3]
Applications Due to its scarcity, high radioactivity, and high toxicity, there are currently no uses for protactinium outside of scientific research. Protactinium-231 (which is formed by the alpha decay of U-235 followed by beta decay of thorium-231) could possibly sustain a nuclear chain reaction. The physicist Walter Seifritz once estimated that protactinium might possibly be used to build a nuclear weapon with a critical mass of 750±180 kg. This possibility (of a chain reaction) has been ruled out by other nuclear physicists since then. The ratio of protactinium-231 to thorium-230 in ocean sediments has also been used in paleoceanography to reconstruct the movements of North Atlantic water bodies during the last melting of Ice Age glaciers.[4]
Protactinium
History In 1890, Mendeleev predicted the existence of an element between thorium and uranium. In 1900, William Crookes isolated protactinium as a radioactive material from uranium; however, he did not identify it as a new element[5] Protactinium was first identified in 1913, when Kasimir Fajans and O. H. Göring encountered the short-lived isotope Pa-234 (half-life of about 1.17 minutes), during their studies of the decay chains of uranium-238 (U-238). They gave the new element the name brevium (from the Latin word, brevis, meaning brief or short);[6] [7] the name was changed to protoactinium in 1918 when two groups of scientists (lead by Otto Hahn and Lise Meitner of Germany; and Frederick Soddy and John Cranston of Great Britain) independently discovered Pa-231. The name was shortened to Protactinium in 1949. Aristid von Grosse prorduced 2 mg of Pa2O5 in 1927,[8] and in 1934 performed the first isolation of elemental protactinium from 0.1 mg of Pa2O5, by converting the oxide to an iodide and then reducing it in a vacuum with an electrically-heated metal filament by the reaction 2PaI5 → 2Pa + 5I2 (iodide process). In 1961, the British Atomic Energy Authority (UKAEA) was able to produce 125 grams of 99.9% pure protactinium by processing 60 tons of waste material in a 12-stage process. For many years, this was the world's only significant supply of protactinium.
Occurrence Protactinium occurs in pitchblende to the extent of about 1.0 part 231Pa per 10 million parts of ore (i.e., 0.1 ppm). Some ores from the Democratic Republic of the Congo have about 3.0 ppm. Protactinium is one of the rarest and most expensive naturally occurring elements.[2]
Compounds Examples of protactinium compounds: • Fluorides: protactinium(IV) fluoride PaF4, protactinium(V) fluoride PaF5 • Chlorides: protactinium(IV) chloride PaCl4, protactinium(V) chloride PaCl5 • Bromides: protactinium(IV) bromide PaBr4, protactinium(V) bromide PaBr5 • Iodides: protactinium(III) iodide PaI3, protactinium(IV) iodide PaI4, protactinium(V) iodide PaI5 • Oxides: protactinium(II) oxide PaO, protactinium(IV) oxide PaO2, protactinium(V) oxide Pa2O5 See also Protactinium compounds.
3
Protactinium
4
Isotopes Twenty-nine radioisotopes of protactinium have been discovered, with the most stable being Pa-231 with a half life of 32760 years, Pa-233 with a half-life of 27.0 days, and Pa-230 with a half-life of 17.4 days. All of the remaining radioactive isotopes have half-lives that are less than 1.60 days, and the majority of these have half-lives that are less than 1.8 seconds. Protactinium also has two meta states, Pa-217m (half-life 1.2 milliseconds) and Pa-234m (half-life 1.17 minutes). The primary decay mode for isotopes of Protactinium lighter than (and including) the most stable isotope Pa-231 (ie, Pa-212 to Pa-231) is alpha decay and the primary mode for the heavier isotopes (ie, Pa-232 to Pa-240) is beta decay. The primary decay products of isotopes of protactinium lighter than (and including) Pa-231 are actinium isotopes and the primary decay products for the heavier isotopes of protactinium are uranium isotopes.
Precautions Protactinium is both toxic and highly radioactive. It requires precautions similar to those used when handling plutonium.
External links • • • •
ANL factsheet [9] WebElements.com - Protactinium It's Elemental - Protactinium [11] PROTACTINIUM [12]
[10]
pnb:مینیٹکیٹورپ
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] C. R. Hammond. The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304857. [3] R. D. Fowler et al. (1965). "Superconductivity of Protactinium". Phys. Rev. Lett. 15: 860. doi: 10.1103/PhysRevLett.15.860 (http:/ / dx. doi. org/ 10. 1103/ PhysRevLett. 15. 860). [4] J. F. McManus, R. Francois, J.-M. Gherardi, L. D. Keigwin, and S. Brown-Leger (2004). "Collapse and rapid resumption of Atlantic meridional circulation linked to deglacial climate changes". Nature 428: 834-837. [5] Emsley, John (2001). " Protactinium (http:/ / books. google. de/ books?id=j-Xu07p3cKwC& pg=PA348)". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. pp. 347–349. ISBN 0198503407. . [6] K. Fajans and 0. Gohring, (1913). " Über die komplexe Natur des Ur X (http:/ / www. digizeitschriften. de/ no_cache/ home/ jkdigitools/ loader/ ?tx_jkDigiTools_pi1[IDDOC]=201162& tx_jkDigiTools_pi1[pp]=425)". Naturwissenschaften 14: 339. doi: 10.1007/BF01495360 (http:/ / dx. doi. org/ 10. 1007/ BF01495360). . [7] K. Fajans and 0. Gohring, (1913). "Über das Uran X2-das neue Element der Uranreihe". Physikalische Zeitschrift 14: 877–84. [8] Aristid von Grosse (1928). "Das Element 91; seine Eigenschaften und seine Gewinnung". Berichte der deutschen chemischen Gesellschaft 61 (1): 233–245. doi: 10.1002/cber.19280610137 (http:/ / dx. doi. org/ 10. 1002/ cber. 19280610137). [9] http:/ / www. ead. anl. gov/ pub/ doc/ protactinium. pdf [10] http:/ / www. webelements. com/ webelements/ elements/ text/ Pa/ index. html [11] http:/ / education. jlab. org/ itselemental/ ele091. html [12] http:/ / pubs. acs. org/ cen/ 80th/ protactinium. html
Article Sources and Contributors
Article Sources and Contributors Protactinium Source: http://en.wikipedia.org/w/index.php?oldid=308866615 Contributors: 65.68.87.xxx, Ahoerstemeier, AlimanRuna, Andres, Archimerged, ArglebargleIV, Arkuat, Bayou Banjo, Benbest, Berkay0652, Bkell, Bobblewik, Bryan Derksen, Chrislk02, Christopher Parham, Conversion script, David Latapie, DerHexer, Dmn, Donarreiskoffer, DragonflySixtyseven, Dwmyers, Edgar181, Egil, Emperorbma, Exert, Femto, Fibonacci, Fresheneesz, Greatpatton, Hak-kâ-ngìn, Hallpriest9, Hammer1980, Hede2000, Heron, Hu, Icairns, Igny, InfoHaunter, Itub, J.delanoy, JWB, Jake Spooky, Jaraalbe, Joanjoc, John, Karelj, Keilana, Kelovy, Kukini, Kwamikagami, LarryMorseDCOhio, Lightscamera808, Marc Venot, Marek69, Materialscientist, Mattman723, Mav, Mgobrien14, Mike Rosoft, Mimihitam, Monkey Bounce, Montgomery '39, NAHID, Nergaal, Nyttend, Olof, Polonium, Poolkris, Pras, Remember, Res2216firestar, Reyk, Romanm, Runt00001, Sam Hocevar, Santăr, Saperaud, Schneelocke, Semperf, Shalom Yechiel, Shoy, Sl, Squids and Chips, Stifynsemons, Stone, Svante, Tagishsimon, Tex, Torsmo, Urhixidur, Usinsk, Vsmith, Vuerqex, Walkerma, Warut, WhiteDragon, Yekrats, Yyy, 107 anonymous edits
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5
Uranium
1
Uranium protactinium ← uranium → neptuniumNd ↑ U ↓ (Uqb)
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
92U Periodic table
Appearance silvery gray metallic; corrodes to a spalling black oxide coat in air
General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Liquid density at m.p.Melting pointBoiling pointHeat of fusionHeat of vaporizationSpecific heat capacityVapor pressure P/Pa
1
10
100
1k
10 k
100 k
Uranium
2 at T/K
2325
2564
2859
3234
3727
4402
Atomic properties Oxidation states ElectronegativityIonization energies 2nd: 1420 kJ·mol−1Atomic radiusCovalent radiusVan der Waals radius Miscellaneous Crystal structureMagnetic orderingElectrical resistivityThermal conductivityThermal expansionSpeed of sound (thin rod) Young's modulusShear modulusBulk modulusPoisson ratioCAS registry number Most stable isotopes Main article: Isotopes of uranium iso
N.A.
half-life
DM
DE (MeV)
DP
232
syn
68.9 y
α & SF
5.414
228
233
syn
159,200 y
SF & α
4.909
229
234
0.0054%
245,500 y
SF & α
4.859
230
235
0.7204%
7.038×108 y
SF & α
4.679
231
236
syn
2.342×107 y
SF & α
4.572
232
238
99.2742%
4.468×109 y
SF & α
4.270
234
U U U U U U
Th Th Th Th Th Th
uranium, U, 92 actiniden/a, 7, f238.02891(3) g·mol−1 [Rn] 5f3 6d1 7s2 2, 8, 18, 32, 21, 9, 2 (Image) solid 19.1 g·cm−3 17.3 g·cm−3 1405.3 K,1132.3 °C,2070 °F 4404 K,4131 °C,7468 °F 9.14 kJ·mol−1 417.1 kJ·mol−1 (25 °C) 27.665 J·mol−1·K−16, 5, 4, 3[1] (weakly basic oxide) 1.38 (Pauling scale) 1st: 597.6 kJ·mol−1156 pm 196±7 pm 186 pm orthorhombic paramagnetic (0 °C) 0.280 µΩ·m (300 K) 27.5 W·m−1·K−1 (25 °C) 13.9 µm·m−1·K−1 (20 °C) 3155 m/s 208 GPa 111 GPa 100 GPa 0.23 7440-61-1 Uranium (pronounced /jʊˈreɪniəm/) is a silvery-white metallic chemical element in the actinide series of the periodic table that has the symbol U and atomic number 92. Besides its 92 protons, a uranium nucleus can have between 141 and 146 neutrons, with 146 (U-238) and 143 (U-235) in its most common isotopes. The number of electrons in a uranium atom is 92, 6 of them valence electrons. Uranium has the highest atomic weight of the naturally occurring elements. Uranium is approximately 70% denser than lead, but not as dense as gold or tungsten. It is weakly radioactive. It occurs naturally in low concentrations (a few parts per million) in soil, rock and water, and is commercially extracted from uranium-bearing minerals such as uraninite (see uranium mining). In nature, uranium atoms exist as uranium-238 (99.284%), uranium-235 (0.711%),[2] and a very small amount of uranium-234 (0.0058%). Uranium decays slowly by emitting an alpha particle. The half-life of uranium-238 is about 4.47 billion years and that of uranium-235 is 704 million years,[3] making them useful in dating the age of the Earth (see uranium-thorium dating, uranium-lead dating and uranium-uranium dating). Many contemporary uses of uranium exploit its unique nuclear properties. Uranium-235 has the distinction of being the only naturally occurring fissile isotope. Uranium-238 is both fissionable by fast neutrons, and fertile (capable of being transmuted to fissile plutonium-239 in a nuclear reactor). An artificial fissile isotope, uranium-233, can be produced from natural thorium and is also important in nuclear technology. While uranium-238 has a small probability to fission spontaneously or when bombarded with fast neutrons, the much higher probability of uranium-235 and to a lesser degree uranium-233 to fission when bombarded with slow neutrons generates the heat in nuclear reactors used as a source of power, and provides the fissile material for nuclear weapons. Both uses rely
Uranium
3
on the ability of uranium to produce a sustained nuclear chain reaction. Depleted uranium (uranium-238) is used in kinetic energy penetrators and armor plating.[4] Uranium is used as a colorant in uranium glass, producing orange-red to lemon yellow hues. It was also used for tinting and shading in early photography. The 1789 discovery of uranium in the mineral pitchblende is credited to Martin Heinrich Klaproth, who named the new element after the planet Uranus. Eugène-Melchior Péligot was the first person to isolate the metal, and its radioactive properties were uncovered in 1896 by Antoine Becquerel. Research by Enrico Fermi and others starting in 1934 led to its use as a fuel in the nuclear power industry and in Little Boy, the first nuclear weapon used in war. An ensuing arms race during the Cold War between the United States and the Soviet Union produced tens of thousands of nuclear weapons that used enriched uranium and uranium-derived plutonium. The security of those weapons and their fissile material following the breakup of the Soviet Union in 1991 is an ongoing concern for public health and safety.
Characteristics When refined, uranium is a silvery white, weakly radioactive metal, which is slightly softer than steel,[5] strongly electropositive and a poor electrical conductor.[6] It is malleable, ductile, and slightly paramagnetic.[5] Uranium metal has very high density, being approximately 70% denser than lead, but slightly less dense than gold. Uranium metal reacts with almost all nonmetallic elements
An induced nuclear fission event involving uranium-235
and their compounds, with reactivity increasing with temperature.[7] Hydrochloric and nitric acids dissolve uranium, but nonoxidizing acids attack the element very slowly.[6] When finely divided, it can react with cold water; in air, uranium metal becomes coated with a dark layer of uranium oxide.[5] Uranium in ores is extracted chemically and converted into uranium dioxide or other chemical forms usable in industry.
Uranium was the first element that was found to be fissile. Upon bombardment with slow neutrons, its uranium-235 isotope will most of the time divide into two smaller nuclei, releasing nuclear binding energy and more neutrons. If these neutrons are absorbed by other uranium-235 nuclei, a nuclear chain reaction occurs and, if there is nothing to absorb some neutrons and slow the reaction, the reaction is explosive. As little as 15 lb (7 kg) of uranium-235 can be used to make an atomic bomb.[8] The first nuclear bomb used in war, Little Boy, relied on uranium fission, while the very first nuclear explosive (The gadget) and the bomb that destroyed Nagasaki (Fat Man) were plutonium bombs. Uranium metal has three allotropic forms:[9] • α (orthorhombic) stable up to 660 °C • β (tetragonal) stable from 660 °C to 760 °C • γ (body-centered cubic) from 760 °C to melting point—this is the most malleable and ductile state.
Uranium
4
Applications Military
Depleted uranium is used by various militaries as high-density penetrators.
The major application of uranium in the military sector is in high-density penetrators. This ammunition consists of depleted uranium (DU) alloyed with 1–2% other elements. At high impact speed, the density, hardness, and flammability of the projectile enable destruction of heavily armored targets. Tank armor and the removable armor on combat vehicles are also hardened with depleted uranium plates. The use of DU became a contentious political-environmental issue after the use of DU munitions by the US, UK and other countries during wars in the Persian Gulf and the Balkans raised questions of uranium compounds left in the soil (see
Gulf War Syndrome).[8] Depleted uranium is also used as a shielding material in some containers used to store and transport radioactive materials.[6] Other uses of DU include counterweights for aircraft control surfaces, as ballast for missile re-entry vehicles and as a shielding material.[5] Due to its high density, this material is found in inertial guidance devices and in gyroscopic compasses.[5] DU is preferred over similarly dense metals due to its ability to be easily machined and cast as well as its relatively low cost.[10] Counter to popular belief, the main risk of exposure to DU is chemical poisoning by uranium oxide rather than radioactivity (uranium being only a weak alpha emitter). During the later stages of World War II, the entire Cold War, and to a lesser extent afterwards, uranium has been used as the fissile explosive material to produce nuclear weapons. Two major types of fission bombs were built: a relatively simple device that uses uranium-235 and a more complicated mechanism that uses uranium-238-derived plutonium-239. Later, a much more complicated and far more powerful fusion bomb that uses a plutonium-based device in a uranium casing to cause a mixture of tritium and deuterium to undergo nuclear fusion was built.[11]
Uranium
5
Civilian The main use of uranium in the civilian sector is to fuel commercial nuclear power plants; by the time it is completely fissioned, one kilogram of uranium-235 can theoretically produce about 20 trillion joules of energy (2 × 1013 joules); as much energy as 1500 tonnes of coal.[4] Commercial nuclear power plants use fuel that is typically enriched to around 3% uranium-235.[4] The CANDU reactor is the only commercial reactor capable of using unenriched uranium fuel. Fuel used for United States Navy reactors is typically highly enriched in uranium-235 (the exact values are classified). In a breeder reactor, uranium-238 can also be converted into plutonium through the following reaction:[5] 238U (n, gamma) → 239U -(beta) → 239Np -(beta) → 239Pu. One
of
the
major
yet-unresolved
issues
with
uranium nuclear fuel is the creation of large amount of nuclear waste. Traditional nuclear reactors burn only 1-2% of uranium fuel. However, it is worth noting that other designs of nuclear reactors using alternative, liquid thorium fuel in molten salt reactors produce virtually no long-lasting nuclear waste.
1oz. sample of U-238 under oil. Surface corrosion is visible on the recently polished surface.
The most visible civilian use of uranium is as the thermal power source used in nuclear power plants.
Uranium
6 Prior to the discovery of radiation, uranium was primarily used in small amounts for yellow glass and pottery glazes (such as uranium glass and in Fiestaware).
After Marie Curie discovered radium in uranium ore, a huge industry developed to mine uranium so as to extract the radium, which was used to make glow-in-the-dark paints for clock and aircraft dials.[12] This left a prodigious quantity of uranium as a 'waste Uranium glass used as lead-in seals in product', since it takes three metric tons of uranium to a vacuum capacitor extract one gram of radium. This 'waste product' was diverted to the glazing industry, making uranium glazes very inexpensive and abundant. In addition to the pottery glazes, uranium tile glazes accounted for the bulk of the use, including common bathroom and kitchen tiles which can be produced in green, yellow, mauve, black, blue, red and other colors. Uranium was also used in photographic chemicals (esp. uranium nitrate as a toner),[5] in lamp filaments, to improve the appearance of dentures, and in the leather and wood industries for stains and dyes. Uranium salts are mordants of silk or wool. Uranyl acetate and uranyl formate are used as electron-dense "stains" in transmission electron microscopy, to increase the contrast of biological specimens in ultrathin sections and in negative staining of viruses, isolated cell organelles and macromolecules.
Uranium glass glowing under UV light
The discovery of the radioactivity of uranium ushered in additional scientific and practical uses of the element. The long half-life of the isotope uranium-238 (4.51 × 109 years) makes it well-suited for use in estimating the age of the earliest igneous rocks and for other types of radiometric dating (including uranium-thorium dating and uranium-lead dating). Uranium metal is used for X-ray targets in the making of high-energy X-rays.[5]
History Pre-discovery use The use of uranium in its natural oxide form dates back to at least the year 79 CE, when it was used to add a yellow color to ceramic glazes.[5] Yellow glass with 1% uranium oxide was found in a Roman villa on Cape Posillipo in the Bay of Naples, Italy by R. T. Gunther of the University of Oxford in 1912.[13] Starting in the late Middle Ages, pitchblende was extracted from the Habsburg silver mines in Joachimsthal, Bohemia (now Jáchymov in the Czech Republic) and was used as a coloring agent in the local glassmaking industry.[14] In the early 19th century, the world's only known sources of uranium ores were these mines.
Uranium
7
Discovery
Antoine Henri Becquerel discovered the phenomenon of radioactivity by exposing a photographic plate to uranium (1896).
The discovery of the element is credited to the German chemist Martin Heinrich Klaproth. While he was working in his experimental laboratory in Berlin in 1789, Klaproth was able to precipitate a yellow compound (likely sodium diuranate) by dissolving pitchblende in nitric acid and neutralizing the solution with sodium hydroxide.[14] Klaproth mistakenly assumed the yellow substance was the oxide of a yet-undiscovered element and heated it with charcoal to obtain a black powder, which he thought was the newly discovered metal itself (in fact, that powder was an oxide of uranium).[14] [15] He named the newly discovered element after the planet Uranus, which had been discovered eight years earlier by William
Herschel.[16] In 1841, Eugène-Melchior Péligot, who was Professor of Analytical Chemistry at the Conservatoire National des Arts et Métiers (Central School of Arts and Manufactures) in Paris, isolated the first sample of uranium metal by heating uranium tetrachloride with potassium.[14] [17] Uranium was not seen as being particularly dangerous during much of the 19th century, leading to the development of various uses for the element. One such use for the oxide was the aforementioned but no longer secret coloring of pottery and glass. Antoine Henri Becquerel discovered radioactivity by using uranium in 1896.[7] Becquerel made the discovery in Paris by leaving a sample of a uranium salt on top of an unexposed photographic plate in a drawer and noting that the plate had become 'fogged'.[18] He determined that a form of invisible light or rays emitted by uranium had exposed the plate.
Fission research A team led by Enrico Fermi in 1934 observed that bombarding uranium with neutrons produces the emission of beta rays (electrons or positrons; see beta particle).[19] The fission products were at first mistaken for new elements of atomic numbers 93 and 94, which the Dean of the Faculty of Rome, Orso Mario Corbino, Enrico Fermi (bottom left) and the rest christened ausonium and hesperium, respectively.[20] of the team that initiated the first [21] [22] [23] The experiments leading to the discovery of artificial nuclear chain reaction (1942). uranium's ability to fission (break apart) into lighter elements and release binding energy were conducted by Otto Hahn and Fritz Strassmann[19] in Hahn's laboratory in Berlin. Lise Meitner and her nephew, physicist Otto Robert Frisch, published the physical explanation in February 1939 and named the process 'nuclear fission'.[24] Soon after, Fermi hypothesized that the fission of uranium might release enough neutrons to sustain a fission reaction. Confirmation of this hypothesis came in 1939, and later work found that on average about 2.5 neutrons are released by each fission of the rare uranium isotope uranium-235.[19] Further work found that the far more common uranium-238 isotope can be transmuted into plutonium, which, like uranium-235,
Uranium
8
is also fissionable by thermal neutrons. These discoveries led numerous countries to begin working on the development of nuclear weapons and nuclear power. On 2 December 1942, as part of the Manhattan Project, another team led by Enrico Fermi was able to initiate the first artificial nuclear chain reaction, Chicago Pile-1. Working in a lab below the stands of Stagg Field at the University of Chicago, the team created the conditions needed for such a reaction by piling together 400 tons (360 tonnes) of graphite, 58 tons (53 tonnes) of uranium oxide, and six tons (five and a half tonnes) of uranium metal.[19]
Bombs
The mushroom cloud over Hiroshima after the dropping of the uranium-based atomic bomb nicknamed 'Little Boy' (1945)
Two major types of atomic bomb were developed by the United States during World War II: a uranium-based device (codenamed "Little Boy") whose fissile material was highly enriched uranium, and a plutonium-based device (see Trinity test and "Fat Man") whose plutonium was derived from uranium-238. The uranium-based Little Boy device became the first nuclear weapon used in war when it was detonated over the Japanese city of Hiroshima on 6 August 1945. Exploding with a yield equivalent to 12,500 tonnes of TNT, the blast and thermal wave of the bomb destroyed nearly 50,000 buildings and killed approximately 75,000 people (see Atomic bombings of Hiroshima and Nagasaki).[18] Initially it was believed that uranium was relatively rare, and that nuclear proliferation could be avoided by simply buying up all known uranium stocks, but within a decade large deposits of it were discovered in many places around the world.[25]
Reactors The X-10 Graphite Reactor at Oak Ridge National Laboratory (ORNL) in Oak Ridge, Tennessee, formerly known as the Clinton Pile and X-10 Pile, was the world's second artificial nuclear reactor (after Enrico Fermi's Chicago Pile) and was the first reactor designed and built for continuous operation. The Experimental Breeder Reactor I at the Idaho National Laboratory (INL) near Arco, Idaho became the first nuclear reactor to create electricity on 20 December 1951. Initially, Four light bulbs lit with electricity four 150-watt light bulbs were lit by the reactor, but generated from the first artificial improvements eventually enabled it to power the whole electricity-producing nuclear reactor, EBR-I (1951) facility (later, the town of Arco became the first in the world to have all its electricity come from nuclear power).[26] The world's first commercial scale nuclear power station, Obninsk in the Soviet
Uranium Union, began generation with its reactor AM-1 on 27 June 1954. Other early nuclear power plants were Calder Hall in England which began generation on 17 October 1956[27] and the Shippingport Atomic Power Station in Pennsylvania which began on 26 May 1958. Nuclear power was used for the first time for propulsion by a submarine, the USS Nautilus, in [19] 1954.
Naturally occurring nuclear fission Fifteen ancient and no longer active natural nuclear fission reactors were found in three separate ore deposits at the Oklo mine in Gabon, West Africa in 1972. Discovered by French physicist Francis Perrin, they are collectively known as the Oklo Fossil Reactors. The ore they exist in is 1.7 billion years old; at that time, uranium-235 constituted about three percent of the total uranium on Earth.[28] This is high enough to permit a sustained nuclear fission chain reaction to occur, providing other conditions are right. The ability of the surrounding sediment to contain the nuclear waste products in less than ideal conditions has been cited by the U.S. federal government as evidence of their claim that the Yucca Mountain nuclear waste repository could safely be a repository of waste for the nuclear power industry.[28]
Cold War legacy and waste During the Cold War between the Soviet Union and the United States, huge stockpiles of uranium were amassed and tens of thousands of nuclear weapons were created using enriched uranium and plutonium made from uranium. Since the break-up of the Soviet Union in 1991, an estimated 600 tons (540 tonnes) of highly enriched weapons grade uranium (enough to make 40,000 nuclear warheads) have been stored in often U.S. and USSR/Russian nuclear weapons stockpiles, 1945–2006 inadequately guarded facilities in the Russian [8] Federation and several other former Soviet states. Police in Asia, Europe, and South America on at least 16 occasions from 1993 to 2005 have intercepted shipments of smuggled bomb-grade uranium or plutonium, most of which was from ex-Soviet sources.[8] From 1993 to 2005 the Material Protection, Control, and Accounting Program, operated by the federal government of the United States, spent approximately US $550 million to help safeguard uranium and plutonium stockpiles in Russia.[8] This money was used for improvements and security enhancements at research and storage facilities. Scientific American reported in February 2006 that some of the facilities security consisted of chain link fences which were in severe states of disrepair. According to an interview from the article, one facility had been storing samples of enriched (weapons grade) uranium in a broom closet prior to the improvement project; another had been keeping track of its stock of nuclear warheads using index cards kept in a shoe box.[29] Above-ground nuclear tests by the Soviet Union and the United States in the 1950s and early 1960s and by France into the 1970s and 1980s[10] spread a significant amount of fallout from uranium daughter isotopes around the world.[30] Additional fallout and pollution occurred from several nuclear accidents.
9
Uranium
10
The Windscale fire at the Sellafield nuclear plant in 1957 spread iodine-131, a short lived radioactive isotope, over much of Northern England. In 1979, the Three Mile Island accident released a small amount of iodine-131. The amounts released by the partial meltdown of the Three Mile Island power plant were minimal, and an environmental survey found only trace amounts in a few field mice dwelling nearby. As I-131 has a half life of slightly more than eight days, any danger posed by the radioactive material has long since passed for both of these incidents. However, the Chernobyl disaster in 1986 was a complete core breach meltdown and partial detonation of the reactor, which ejected iodine-131 and strontium-90 over a large area of Europe. The 28 year half-life of strontium-90 has only recently allowed some of the surrounding countryside around the reactor to be habitable.[10] Since this is less than one half life after the accident, more than half of the original release of strontium-90 will still be present.
Occurrence Biotic and abiotic Uranium is a naturally occurring element that can be
Uraninite, also known as Pitchblende, is the most common ore mined to extract uranium.
found in low levels within all rock, soil, and water. Uranium is also the highest-numbered element to be found naturally in significant quantities on earth and is always found combined with other elements.[5] Along with all elements having atomic weights higher than that of iron, it is only naturally formed in supernovas.[31] The decay of uranium, thorium, and potassium-40 in the Earth's mantle is thought to be the main source of heat[32] [33] that keeps the outer core liquid and drives mantle convection, which in turn drives plate tectonics.
Uranium's average concentration in the Earth's crust is (depending on the reference) 2 to 4 parts per million,[6] [10] or about 40 times as abundant as silver.[7] The Earth's crust from the surface to 25 km (15 mi) down is calculated to contain 1017 kg (2 × 1017 lb) of uranium while the oceans may contain 1013 kg (2 × 1013 lb).[6] The concentration of uranium in soil ranges from 0.7 to 11 parts per million (up to 15 parts per million in farmland soil due to use of phosphate fertilizers), and its concentration in sea water is 3 parts per billion.[10] Uranium is more plentiful than antimony, tin, cadmium, mercury, or silver, and it is about as abundant as arsenic or molybdenum.[5] [10] Uranium is found in hundreds of minerals including uraninite (the most common uranium ore), carnotite, autunite, uranophane, torbernite, and coffinite.[5] Significant concentrations of uranium occur in some substances such as phosphate rock deposits, and minerals such as lignite, and monazite sands in uranium-rich ores[5] (it is recovered commercially from sources with as little as 0.1% uranium[7] ).
Uranium
Some organisms, such as the lichen Trapelia involuta or microorganisms such as the bacterium Citrobacter, can absorb concentrations of uranium that are up to 300 times higher than in their environment.[34] Citrobacter species absorb uranyl ions when given glycerol phosphate (or other similar organic phosphates). After one day, one gram of bacteria can encrust themselves with nine grams of uranyl phosphate crystals; this creates the possibility that these organisms could be used in bioremediation to decontaminate uranium-polluted water.[14] [35]
11
Citrobacter species can have concentrations of uranium in their bodies 300 times higher than in the surrounding environment.
In nature, uranium (VI) forms highly soluble carbonate complexes at alkaline pH. This leads to an increase in mobility and availability of uranium to groundwater and soil from nuclear wastes which leads to health hazards. However, it is difficult to precipitate uranium as phosphate in the presence of excess carbonate at alkaline pH. A Sphingomonas sp. strain BSAR-1 has been found to express a high activity alkaline phosphatase (PhoK) that has been applied for bioprecipitation of uranium as uranyl phosphate species from alkaline solutions. The precipitation ability was enhanced by overexpressing PhoK protein in E. coli.[36] Plants absorb some uranium from soil. Dry weight concentrations of uranium in plants range from 5 to 60 parts per billion, and ash from burnt wood can have concentrations up to 4 parts per million.[14] Dry weight concentrations of uranium in food plants are typically lower with one to two micrograms per day ingested through the food people eat.[14]
Production and mining The worldwide production of uranium in 2006 amounted to 39 655 tonnes, of which 25% was mined in Canada. Other important uranium mining countries are Australia (19.1%), Kazakhstan (13.3%), Niger (8.7%), Russia (8.6%), and Namibia (7.8%). Uranium ore is mined in several ways: by open pit, underground, in-situ leaching, and borehole mining (see uranium mining).[4] Low-grade uranium ore mined in 2006 typically contains 0.01 to 0.25% uranium oxides. Extensive measures must be employed to extract the metal from its ore.[37] High-grade ores found in Athabasca Basin deposits in Saskatchewan, Canada can Yellowcake is a concentrated mixture contain up to 70% uranium oxides, and therefore must of uranium oxides that is further be diluted with rock prior to milling, in order to reduce refined to extract pure uranium. radiation exposure to workers. Uranium ore is crushed and rendered into a fine powder and then leached with either an acid or alkali. The leachate is subjected to one of several sequences of precipitation, solvent extraction, and ion exchange. The resulting mixture, called yellowcake, contains at least 75% uranium oxides. Yellowcake is then calcined to remove impurities from the milling process prior to refining and conversion. Commercial-grade uranium can be produced through the reduction of uranium halides with alkali or alkaline earth metals.[5] Uranium metal can also be made through electrolysis of
Uranium KU5 or UF4, dissolved in molten calcium chloride (CaCl2) and sodium chloride (NaCl) solution.[5] Very pure uranium can be produced through the thermal decomposition of uranium halides on a hot filament.[5] Miners who worked in uranium filled mines have a very high incidence of cancer relative to the rest of the United States population. Though the Navajo workers and families noticed this in the 1950s, bureaucrats dragged their feet, and companies disregarded warnings. The miners, especially the Navajo miners, were kept from receiving compensation for the suffering they went through. In 1990 a law was passed known as the Radiation Exposure Compensation Act of 1990 (RECA) (Eichstaedt, 1994). The law required $100,000 in "compassion payments" to uranium miners diagnosed with cancer or other respiratory ailments (Eichstaedt, 1994; Benally Sr., 1995). To qualify for compensation, a miner had to prove that s/he had worked in the mines and was now suffering from one of the diseases on the compensation list (Eichstaedt, 1994; Benally Sr. 1995).[38]
Resources and reserves Current economic uranium resources will last for over 100 years at 2006 consumption rates, while it is expected there is twice that amount awaiting discovery. With reprocessing and recycling, the reserves are good for thousands of years.[39] It is estimated that 5.5 million tonnes of uranium ore reserves are economically viable at US$59/lb,[39] while 35 million tonnes are classed as mineral resources (reasonable prospects for eventual economic extraction).[40] An additional 4.6 billion tonnes of uranium are estimated to be in sea water (Japanese scientists in the 1980s showed that extraction of uranium from sea water using ion exchangers was feasible).[41] [42] Exploration for uranium is increasing with US$200 million being spent world wide in 2005, a 54% increase on the previous year.[40] This trend continued through 2006, when expenditure on exploration rocketed to over $774 million, an increase of over 250% compared to 2004. The OECD Nuclear Energy Agency said exploration figures for 2007 would likely match those for 2006.[39] Australia has 23% of the world's uranium ore reserves[43] and the world's largest single uranium deposit, located at the Olympic Dam Mine in South Australia.[44] Almost all Australia's mined uranium is exported, under strict International Atomic Energy Agency safeguards against use in nuclear weapons. Some nuclear fuel comes from nuclear weapons being dismantled.[45]
Supply In 2005, seventeen countries produced concentrated uranium oxides, with Canada (27.9% of world production) and Australia (22.8%) being the largest producers and Kazakhstan (10.5%), Russia (8.0%), Namibia (7.5%), Niger (7.4%), Uzbekistan (5.5%), the United States (2.5%), Argentina (2.1%), Ukraine (1.9%) Uranium output in 2005 and China (1.7%) also producing significant amounts.[46] Kazakhstan continues to increase production and may become the world's largest producer of uranium by this year (2009) with an expected production of
12
Uranium
13
12,826 tonnes, compared to Canada with 11,100 tonnes and Australia with 9,430 tonnes.[47] [48] The ultimate available uranium is believed to be sufficient for at least the next 85 years[40] although some studies indicate underinvestment in the late twentieth century may produce supply problems in the 21st century.[49] Some claim that production of uranium will peak similar to peak oil. Kenneth S. Deffeyes and Ian D. MacGregor point out that uranium deposits seem to be log-normal distributed. There is a 300-fold increase in the amount of uranium recoverable for each tenfold decrease in ore grade."[50] In other words, there is little high grade ore and proportionately much more low grade ore available.
Compounds Oxidation states and oxides Oxides Calcined uranium yellowcake as produced in many large mills contains a distribution of uranium oxidation species in various forms ranging from most oxidized to least oxidized. Particles with short residence times in a calciner will generally be less oxidized than those with long retention times or particles recovered in the stack scrubber. Uranium content is usually referenced to U3O8, which dates to the days of the Manhattan project when U3O8 was used as an analytical chemistry reporting standard.
Triuranium octaoxide (diagram pictured) and uranium dioxide are the two most common uranium oxides.
Phase relationships in the uranium-oxygen system are complex. The most important oxidation states of uranium are uranium(IV) and uranium(VI), and their two corresponding oxides are, respectively, uranium dioxide (UO2) and uranium trioxide (UO3).[51] Other uranium oxides such as uranium monoxide (UO), diuranium pentoxide (U2O5), and uranium peroxide (UO4•2H2O) also exist. The most common forms of uranium oxide are triuranium octaoxide (U3O8) and UO2.[52] Both oxide forms are solids that have low solubility in water and are relatively stable over a wide range of environmental conditions. Triuranium octaoxide is (depending on conditions) the most stable compound of uranium and is the form most commonly found in nature. Uranium dioxide is the form in which uranium is most commonly used as a nuclear reactor fuel.[52] At ambient temperatures, UO2 will gradually convert to U3O8. Because of their stability, uranium oxides are generally considered the preferred chemical form for storage or disposal.[52] Aqueous chemistry The four different oxidation states of uranium are soluble and therefore can be studied in aqueous solutions. They are: U3+ (red), U4+ (green), UO2+ (unstable), and UO22+ (yellow).[53] A few solid and semi-metallic compounds such as UO and US exist for the formal oxidation state uranium(II), but no simple ions are known to exist in solution for that state. Ions of U3+ liberate hydrogen from water and are therefore considered to be highly unstable. The UO22+ ion represents the uranium(VI) state and is known to form compounds
Uranium
14
such as carbonate, chloride and sulfate. UO22+ also forms complexes with various organic chelating agents, the most commonly encountered of which is uranyl acetate.[53] Carbonates The interactions of carbonate anions with uranium(VI) cause the Pourbaix diagram to change greatly when the medium is changed from water to a carbonate containing solution. It is interesting to note that while the vast majority of carbonates are insoluble in water (students are often taught that all carbonates other than those of alkali metals are insoluble in water), uranium carbonates are often soluble in water. This is due to the fact that a U(VI) cation is able to bind two terminal oxides and three or more carbonates to form anionic complexes.
The Pourbaix diagram for uranium in a non-complexing aqueous medium (eg perchloric [54] acid / sodium hydroxide).
The Pourbaix diagram for uranium in carbonate [54] solution
Uranium
15
The effect of pH The uranium fraction diagrams in the presence of carbonate illustrate this further: when the pH of a uranium(VI) solution increases, the uranium is converted to a hydrated uranium oxide hydroxide and at high pHs it becomes an anionic hydroxide complex.
A diagram showing the relative concentrations of the different chemical forms of uranium in a non-complexing aqueous medium (eg perchloric [54] acid / sodium hydroxide).
When carbonate is added, uranium is converted to a series of carbonate complexes if the pH is increased. One effect of these reactions is increased solubility of uranium in the pH range 6 to 8, a fact which has a direct bearing on the long term stability of spent uranium dioxide nuclear fuels.
Hydrides, carbides and nitrides Uranium metal heated to 250 to 300 °C (482
A diagram showing the relative concentrations of the different chemical forms of uranium in an [54] aqueous carbonate solution.
to 572 °F) reacts with hydrogen to form uranium hydride. Even higher temperatures will reversibly remove the hydrogen. This property makes uranium hydrides convenient starting materials to create reactive uranium powder along with various uranium carbide, nitride, and halide compounds.[55] Two crystal modifications of uranium hydride exist: an α form that is obtained at low temperatures and a β form that is created when the formation temperature is above 250 °C.[55]
Uranium carbides and uranium nitrides are both relatively inert semimetallic compounds that are minimally soluble in acids, react with water, and can ignite in air to form U3O8.[55] Carbides of uranium include uranium monocarbide (UC), uranium dicarbide (UC2), and diuranium tricarbide (U2C3). Both UC and UC2 are formed by adding carbon to molten uranium or by exposing the metal to carbon monoxide at high temperatures. Stable below 1800 °C, U2C3 is prepared by subjecting a heated mixture of UC and UC2 to mechanical stress.[56] Uranium nitrides obtained by direct exposure of the metal to nitrogen include uranium mononitride (UN), uranium dinitride (UN2), and diuranium trinitride (U2N3).[56]
Uranium
16
Halides All uranium fluorides are created using uranium tetrafluoride (UF4); UF4 itself is prepared by hydrofluorination of uranium dioxide.[55] Reduction of UF4 with hydrogen at 1000 °C produces uranium trifluoride (UF3). Under the right conditions of temperature and pressure, the reaction of solid UF4 with gaseous uranium hexafluoride (UF6) can form the intermediate fluorides of U2F9, U4F17, and UF5.[55] At room temperatures, UF6 has a high vapor pressure, making it useful in the gaseous diffusion process to separate uranium-235 from the common uranium-238 isotope. This compound can be prepared from uranium dioxide and uranium hydride by the following process:[55] UO2 + 4 HF → UF4 + 2 H2O (500 °C, endothermic)
Uranium hexafluoride is the feedstock used to separate uranium-235 from natural uranium.
UF4 + F2 → UF6 (350 °C, endothermic) The resulting UF6, a white solid, is highly reactive (by fluorination), easily sublimes (emitting a nearly perfect gas vapor), and is the most volatile compound of uranium known to exist.[55] One method of preparing uranium tetrachloride (UCl4) is to directly combine chlorine with either uranium metal or uranium hydride. The reduction of UCl4 by hydrogen produces uranium trichloride (UCl3) while the higher chlorides of uranium are prepared by reaction with additional chlorine.[55] All uranium chlorides react with water and air. Bromides and iodides of uranium are formed by direct reaction of, respectively, bromine and iodine with uranium or by adding UH3 to those element's acids.[55] Known examples include: UBr3, UBr4, UI3, and UI4. Uranium oxyhalides are water-soluble and include UO2F2, UOCl2, UO2Cl2, and UO2Br2. Stability of the oxyhalides decrease as the atomic weight of the component halide increases.[55]
Isotopes Natural concentrations Natural uranium consists of three major isotopes: uranium-238 (99.28% natural abundance), uranium-235 (0.71%), and uranium-234 (0.0054%). All three are radioactive. Uranium-238 is the most stable isotope, with a half-life of 4.51 × 109 years (close to the age of the Earth). Uranium-235 has a half-life of 7.13 × 108 years, and uranium-234 has a half-life of 2.48 × 105 years.[57] Uranium-238 is an α emitter, decaying through the 18-member uranium natural decay series into lead-206.[7] The decay series of uranium-235 (also called actino-uranium) has 15 members that ends in lead-207.[7] The constant rates of decay in these series makes comparison of the ratios of parent to daughter elements useful in radiometric dating. Uranium-234 decays to lead-206 through a series of short-lived intermediaries. Uranium-233 is made from thorium-232 by neutron bombardment;[5] its decay series ends
Uranium
17
with thallium-205. The isotope uranium-235 is important for both nuclear reactors and nuclear weapons because it is the only isotope existing in nature to any appreciable extent that is fissile, that is, can be broken apart by thermal neutrons.[7] The isotope uranium-238 is also important because it absorbs neutrons to produce a radioactive isotope that subsequently decays to the isotope plutonium-239, which is also fissile.[19]
Enrichment Isotope separation concentrates (enriches) the fissionable uranium-235 for nuclear weapons and most nuclear power plants, with the exception of gas cooled reactors and pressurised heavy water reactors. Most neutrons released by a fissioning atom of uranium-235 must impact other uranium-235 atoms to sustain the nuclear chain reaction. The concentration and amount of uranium-235 needed to achieve this is called a 'critical mass'. Cascades of gas centrifuges are used to enrich uranium ore to concentrate its fissionable isotopes.
To be considered 'enriched', the uranium-235 fraction should be between 3% and 5%.[58] This process produces huge quantities of uranium that is depleted of uranium-235 and with a correspondingly increased fraction of uranium-238, called depleted uranium or 'DU'. To be considered 'depleted', the uranium-235 isotope concentration should be no more than 0.2% to 0.3%.[59] The price of uranium has risen since 2001, so enrichment tailings containing more than 0.35% uranium-235 are being considered for re-enrichment, driving the price of depleted uranium hexafluoride above $130 per kilogram in July, 2007 from $5 in 2001.[59] The gas centrifuge process, where gaseous uranium hexafluoride (UF6) is separated by the difference in molecular weight between 235UF6 and 238UF6 using high-speed centrifuges, is [18] the cheapest and leading enrichment process. The gaseous diffusion process had been the leading method for enrichment and was used in the Manhattan Project. In this process, uranium hexafluoride is repeatedly diffused through a silver-zinc membrane, and the different isotopes of uranium are separated by diffusion rate (since uranium 238 is heavier it diffuses slightly slower than uranium-235).[18] The molecular laser isotope separation method employs a laser beam of precise energy to sever the bond between uranium-235 and fluorine. This leaves uranium-238 bonded to fluorine and allows uranium-235 metal to precipitate from the solution.[4] Another method used is liquid thermal diffusion.[6]
Precautions Exposure A person can be exposed to uranium (or its radioactive daughters such as radon) by inhaling dust in air or by ingesting contaminated water and food. The amount of uranium in air is usually very small; however, people who work in factories that process phosphate fertilizers, live near government facilities that made or tested nuclear weapons, live or work near a modern battlefield where depleted uranium weapons have been used, or live or work
Uranium
18
near a coal-fired power plant, facilities that mine or process uranium ore, or enrich uranium for reactor fuel, may have increased exposure to uranium.[60] [61] Houses or structures that are over uranium deposits (either natural or man-made slag deposits) may have an increased incidence of exposure to radon gas. Almost all uranium that is ingested is excreted during digestion, but up to 5% is absorbed by the body when the soluble uranyl ion is ingested while only 0.5% is absorbed when insoluble forms of uranium, such as its oxide, are ingested.[14] However, soluble uranium compounds tend to quickly pass through the body whereas insoluble uranium compounds, especially when ingested via dust into the lungs, pose a more serious exposure hazard. After entering the bloodstream, the absorbed uranium tends to bioaccumulate and stay for many years in bone tissue because of uranium's affinity for phosphates.[14] Uranium is not absorbed through the skin, and alpha particles released by uranium cannot penetrate the skin.
Effects Normal functioning of the kidney, brain, liver, heart, and other systems can be affected by uranium exposure, because, in addition to being weakly radioactive, uranium is a toxic metal.[14] [62] [63] Uranium is also a reproductive toxicant.[64] [65] Radiological effects are generally local because alpha radiation, the primary form of U-238 decay, has a very short range, and will not penetrate skin. Uranyl (UO2+) ions, such as from uranium trioxide or uranyl nitrate and other hexavalent uranium compounds, have been shown to cause birth defects and immune system damage in laboratory animals.[66] While the CDC has published one study that no human cancer has been seen as a result of exposure to natural or depleted uranium,[67] exposure to uranium and its decay products, especially radon, are widely known and significant health threats.[10] Exposure to strontium-90, iodine-131, and other fission products is unrelated to uranium exposure, but may result from medical procedures or exposure to spent reactor fuel or fallout from nuclear weapons.[68] Although accidental inhalation exposure to a high concentration of uranium hexafluoride has resulted in human fatalities, those deaths were not associated with uranium itself.[69] Finely divided uranium metal presents a fire hazard because uranium is pyrophoric; small grains will ignite spontaneously in air at room temperature.[5] Compilation Body system
[70]
[62] of 2004 Review Information Regarding Uranium Toxicity
Human studies
Animal studies
In vitro
Renal
Elevated levels of protein excretion, urinary catalase and diuresis
Damage to Proximal convoluted tubules, necrotic cells cast from tubular epithelium, glomerular changes
No studies
Brain/CNS
Decreased performance on neurocognitive tests
Acute cholinergic toxicity; Dose-dependent accumulation in cortex, midbrain, and vermis; Electrophysiological changes in hippocampus
No studies
Uranium
19
DNA
Increased reports of cancers
Increased urine mutagenicity and induction of tumors
Binucleated cells with micronuclei, Inhibition of cell cycle kinetics and proliferation; Sister chromatid induction, tumorigenic phenotype
Bone/muscle
No studies
Inhibition of periodontal bone formation; and alveolar wound healing
No studies
Reproductive
Uranium miners have more first born female children
Moderate to severe focal tubular atrophy; vacuolization of Leydig cells
No studies
Lungs/respiratory
No adverse health effects reported
Severe nasal congestion and hemorrage, lung lesions and fibrosis, edema and swelling, lung cancer
No studies
Gastrointestinal
Vomiting, diarrhea, albuminuria
n/a
n/a
Liver
No effects seen at exposure dose
Fatty livers, focal necrosis
No studies
Skin
No exposure assessment data available
Swollen vacuolated epidermal cells, No studies damage to hair follicles and sebaceous glands
Tissues surrounding embedded DU fragments
Elevated uranium urine concentrations
Elevated uranium urine concentrations, perturbations in biochemical and neuropsychological testing
No studies
Immune system
Chronic fatigue, rash, ear and eye infections, hair and weight loss, cough. May be due to combined chemical exposure rather than DU alone
No studies
No studies
Eyes
No studies
Conjunctivitis, irritation inflammation, edema, ulceration of conjunctival sacs
No studies
Blood
No studies
Decrease in RBC count and hemoglobin concentration
No studies
Cardiovascular
Myocarditis resulting from No effects the uranium ingestion, which ended 6 months after ingestion
No studies
Uranium
See also • • • • • • • • • • • • •
Nuclear engineering Nuclear fuel cycle Thorium Thorium fuel cycle Plutonium Nuclear physics K-65 residues List of uranium mines Isotopes of uranium Uranate, an anion of uranium Uranium leak Uranium glass Uranium reserves
References Full reference information for multi-page works cited • John Emsley (2001). "Uranium [71]". Nature's Building Blocks: An A to Z Guide to the Elements. Oxford: Oxford University Press. pp. 476–82. ISBN 0-19-850340-7. http:/ / books. google. com/ books?id=j-Xu07p3cKwC& printsec=frontcover. • Glenn T. Seaborg (1968). "Uranium". The Encyclopedia of the Chemical Elements. Skokie, Illinois: Reinhold Book Corporation. pp. 773–86. LCCCN 68-29938.
External links • ATSDR Case Studies in Environmental Medicine: Uranium Toxicity [72] U.S. Department of Health and Human Services • "Public Health Statement for Uranium [73]". CDC. http:/ / www. atsdr. cdc. gov/ toxprofiles/ phs150. html. • Uranium Resources and Nuclear Energy [74] • U.S. EPA: Radiation Information for Uranium [75] • "What is Uranium?" from World Nuclear Association [76] • Nuclear fuel data and analysis from the U.S. Energy Information Administration [77] • Current market price of uranium [78] • World Uranium deposit maps [79] • Annotated bibliography for uranium from the Alsos Digital Library [80] • NLM Hazardous Substances Databank—Uranium, Radioactive [81] • 'Pac-Man' molecule chews up uranium contamination - earth - 17 January 2008 - New Scientist Environment [82] • Mining Uranium at Namibia's Langer Heinrich Mine [83] • Uranium futures market [84] • World Nuclear News [85] ckb:مویناروی
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Uranium
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Uranium S0012-821X(02)00930-5). [31] " WorldBook@NASA: Supernova (http:/ / www. nasa. gov/ worldbook/ supernova_worldbook_prt. htm)". NASA. . Retrieved 2007-02-19. [32] Biever, Celeste (27 July 2005). First measurements of Earth's core radioactivity (http:/ / www. newscientist. com/ channel/ earth/ mg18725103. 700). New Scientist. . [33] " Potassium-40 heats up Earth's core (http:/ / physicsworld. com/ cws/ article/ news/ 17436)". physicsweb. 7 May 2003. . Retrieved 2007-01-14. [34] Emsley, pages 476 and 482 [35] L. E. Macaskie, R. M. Empson, A. K. Cheetham, C. P. Grey, A. J. Skarnulis (1992). "Uranium bioaccumulation by a Citrobacter sp. as a result of enzymically mediated growth of polycrystalline HUO2PO4". Science 257: 782–784. doi: 10.1126/science.1496397 (http:/ / dx. doi. org/ 10. 1126/ science. 1496397). PMID 1496397. [36] K.S. Nilgiriwala, A. Alahari, A. S. Rao & S.K. Apte (September 2008). "Cloning and overexpression of an alkaline phosphatase PhoK from Sphingomonas sp. BSAR-1 for uranium bioprecipitation from alkaline solutions". Applied and Environmental Microbiology 74 (17): 5516-5523. doi: 10.1128/AEM.00107-08 (http:/ / dx. doi. org/ 10. 1128/ AEM. 00107-08). ISSN 1098-5336 (http:/ / worldcat. org/ issn/ 1098-5336). PMID 18641147. [37] Seaborg, page 774 [38] " Environmental Justice for the Navajo : Uranium Mining in the Southwest (http:/ / www. umich. edu/ ~snre492/ sdancy. html)". University of Michigan. . Retrieved 2009-08-08. [39] " Exploration drives uranium resources up 17% (http:/ / www. world-nuclear-news. org/ ENF_Exploration_drives_uranium_resources_up_17_0206082. html)". World-nuclear-news.org. . Retrieved 2008-09-12. [40] " Global Uranium Resources to Meet Projected Demand (http:/ / www. iaea. org/ NewsCenter/ News/ 2006/ uranium_resources. html)". International Atomic Energy Agency. 2006. . Retrieved 2007-03-29. [41] " Uranium recovery from Seawater (http:/ / www. jaea. go. jp/ jaeri/ english/ ff/ ff43/ topics. html)". Japan Atomic Energy Research Institute. 1999-08-23. . Retrieved 2008-09-03. [42] " How long will nuclear energy last? (http:/ / www-formal. stanford. edu/ jmc/ progress/ cohen. html)". 1996-02-12. . Retrieved 2007-03-29. [43] " Supply of Uranium (http:/ / www. world-nuclear. org/ info/ inf75. html)". . [44] " Uranium Mining and Processing in South Australia (http:/ / www. uraniumsa. org/ processing/ processing. htm)". South Australian Chamber of Mines and Energy. 2002. . Retrieved 2007-01-14. [45] Military Warheads as a Source of Nuclear Fuel (http:/ / www. world-nuclear. org/ info/ inf13. html) [46] " World Uranium Production (http:/ / www. uxc. com/ fuelcycle/ uranium/ production-uranium. html)". UxC Consulting Company, LLC. . Retrieved 2007-02-11. [47] Posted by Mithridates (July 24, 2008). " Page F30: Kazakhstan to surpass Canada as the world's largest producer of uranium by next year (2009) (http:/ / mithridates. blogspot. com/ 2008/ 07/ kazakhstan-to-surpass-canada-as-worlds. html)". Mithridates.blogspot.com. . Retrieved 2008-09-12. [48] " ZAMAN GAZETESİ [İnternetin İlk Türk Gazetesi] - Kazakistan uranyum üretimini artıracak (http:/ / www. zaman. com. tr/ haber. do?haberno=717292)" (in tr). Zaman.com.tr. . Retrieved 2008-09-12. [49] " Lack of fuel may limit U.S. nuclear power expansion (http:/ / web. mit. edu/ newsoffice/ 2007/ fuel-supply. html)". Massachusetts Institute of Technology. 2007-03-21. . Retrieved 2007-03-29. [50] Kenneth S. Deffeyes and Ian D. MacGregor (1980-01). " World Uranium Resources (http:/ / www. osti. gov/ energycitations/ product. biblio. jsp?osti_id=6665051)". Scientific American. p. p 66. . Retrieved 2008-04-21. [51] Seaborg, page 779 [52] " Chemical Forms of Uranium (http:/ / web. ead. anl. gov/ uranium/ guide/ ucompound/ forms/ index. cfm)". Argonne National Laboratory. . Retrieved 2007-02-18. [53] Seaborg, page 778 [54] Ignasi Puigdomenech, Hydra/Medusa Chemical Equilibrium Database and Plotting Software (2004) KTH Royal Institute of Technology, freely downloadable software at (http:/ / www. kemi. kth. se/ medusa/ ) [55] Seaborg, page 782 [56] Seaborg, page 780 [57] Seaborg, page 777 [58] " Uranium Enrichment (http:/ / web. ead. anl. gov/ uranium/ guide/ depletedu/ enrich/ index. cfm)". Argonne National Laboratory. . Retrieved 2007-02-11. [59] Peter Diehl. " Depleted Uranium: a by-product of the Nuclear Chain (http:/ / www. wise-uranium. org/ dhap991. html)". Laka Foundation. . Retrieved 2009-07-31. [60] " Radiation Information for Uranium (http:/ / www. epa. gov/ radiation/ radionuclides/ uranium. html)". U.S. Environmental Protection Agency. . Retrieved 2009-07-31.
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Uranium [61] " ToxFAQ for Uranium (http:/ / www. atsdr. cdc. gov/ tfacts150. html)". Agency for Toxic Substances and Disease Registry. September 1999. . Retrieved 2007-02-18. [62] E. S. Craft, A. W. Abu-Qare, M. M. Flaherty, M. C. Garofolo, H. L. Rincavage, M. B. Abou-Donia (2004). "Depleted and natural uranium: chemistry and toxicological effects". Journal of Toxicology and Environmental Health Part B: Critical Reviews 7 (4): 297–317. doi: 10.1080/10937400490452714 (http:/ / dx. doi. org/ 10. 1080/ 10937400490452714). [63] " Toxicological Profile for Uranium (http:/ / www. atsdr. cdc. gov/ toxprofiles/ tp150-c2. pdf)" (PDF). Atlanta, GA: Agency for Toxic Substances and Disease Registry (ATSDR). 1999. . [64] Hindin, et al. (2005) "Teratogenicity of depleted uranium aerosols: A review from an epidemiological perspective," (http:/ / www. pubmedcentral. nih. gov/ articlerender. fcgi?artid=1242351) Environ Health, vol. 4, pp. 17 [65] Arfsten, D.P.; K.R. Still; G.D. Ritchie (2001). "A review of the effects of uranium and depleted uranium exposure on reproduction and fetal development". Toxicology and Industrial Health 17 (5–10): 180–91. doi: 10.1191/0748233701th111oa (http:/ / dx. doi. org/ 10. 1191/ 0748233701th111oa). PMID 12539863. [66] Domingo, J. (2001) "Reproductive and developmental toxicity of natural and depleted uranium: a review," Reproductive Toxicology, vol. 15, pp. 603–609, doi: 10.1016/S0890-6238(01)00181-2 PMID 2711400 [67] " Public Health Statement for Uranium (http:/ / www. atsdr. cdc. gov/ toxprofiles/ phs150. html)". CDC. . Retrieved 2007-02-15. [68] Chart of the Nuclides, US Atomic Energy Commission 1968 [69] Kathren and Moore 1986; Moore and Kathren 1985; USNRC 1986 [70] http:/ / www. informaworld. com/ smpp/ section?content=a714111273& fulltext=713240929#uteb_12817_tab07 [71] http:/ / books. google. com/ books?id=j-Xu07p3cKwC& printsec=frontcover [72] http:/ / www. atsdr. cdc. gov/ csem/ uranium/ [73] http:/ / www. atsdr. cdc. gov/ toxprofiles/ phs150. html [74] http:/ / www. lbst. de/ publications/ studies__e/ 2006/ EWG-paper_1-06_Uranium-Resources-Nuclear-Energy_03DEC2006. pdf [75] http:/ / www. epa. gov/ radiation/ radionuclides/ uranium. html [76] http:/ / world-nuclear. org/ education/ uran. htm [77] http:/ / www. eia. doe. gov/ fuelnuclear. html [78] http:/ / www. uxc. com [79] http:/ / www. antenna. nl/ wise/ uranium/ umaps. html [80] http:/ / alsos. wlu. edu/ qsearch. aspx?browse=science/ Uranium [81] http:/ / toxnet. nlm. nih. gov/ cgi-bin/ sis/ search/ r?dbs+ hsdb:@term+ @na+ @rel+ uranium,+ radioactive [82] http:/ / environment. newscientist. com/ article/ mg19726396. 200-pacman-molecule-chews-up-uranium-contamination. html [83] http:/ / viewer. zmags. com/ showmag. php?mid=pfgsh#/ page34/ [84] http:/ / www. nymex. com/ UX_pre_agree. aspx [85] http:/ / www. world-nuclear-news. org/
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Article Sources and Contributors
Article Sources and Contributors Uranium Source: http://en.wikipedia.org/w/index.php?oldid=308226780 Contributors: .:Ajvol:., .:CoReHaCk:., 10strein, 13afuse, 144.132.70.xxx, 200.191.188.xxx, 2D, 3sXk3, 56, A.Ou, ABF, ASDFGHJKL, AVand, Abopardikar, Abrech, Aces lead, Aciel, Acroterion, Adashiel, Aff123a, Ahoerstemeier, Airbreather, Aitias, Alansohn, AlexiusHoratius, Alexparent, Alhutch, AlmostReadytoFly, Alphachimp, Altus Prosator, Amarkov, Amonkey person123456, Amplitude101, Anaxial, Andesite, AndonicO, Andres, Andrewa, Angela, Angusmclellan, Anlace, Anomalocaris, Antandrus, Antatnsu, Antonio Lopez, Anwar saadat, Apollo2011, Apotheosis247, Arakunem, Archimerged, Arenhaus, Arjun01, ArnoldReinhold, Art10, Artaxiad, Artichoker, Artivist, Ashley Pomeroy, Aspects, Asshairbutt, Atarr, Atomicskier, Atoyotaisatoyota, Austinfidel, Awadewit, Axl, AySz88, B.d.mills, BCV, Babajobu, Barneca, Barneyg, Barticus88, Baum, Bayou Banjo, Bbatsell, Bduke, Beagle17, Beast of traal, BeefRendang, Beetstra, 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alf, Wiki ian, WikiTed 2000, Wikibofh, Wikidudeman, Wimt, Wimvandorst, Wisdom89, Witoman, Wknight94, Woseph, Wowlookitsjoe, Wtmitchell, WulfTheSaxon, X201, Xiahou, Xxpor, Xy7, Yann, Yapete, Yekrats, Yellowdesk, Yggdræsil, Yonatan, Ytterbium2, Yyy, Zapvet, Zeno Panthakree, Zeszezu, Zscout370, Zzuuzz, Zzyzx11, دمحا-بيجن-ينابايب-لېخميهاربا, 1624 anonymous edits
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Image Sources, Licenses and Contributors Image:Vacuum capacitor with uranium glass.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Vacuum_capacitor_with_uranium_glass.jpg License: unknown Contributors: User:Warut Image:U glass with black light.jpg Source: http://en.wikipedia.org/w/index.php?title=File:U_glass_with_black_light.jpg License: Creative Commons Attribution-Sharealike 2.5 Contributors: Nolanus, Noodle snacks, Túrelio, Zvesoulis Image:Becquerel plate.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Becquerel_plate.jpg License: Public Domain Contributors: Ranveig Image:ChicagoPileTeam.png Source: http://en.wikipedia.org/w/index.php?title=File:ChicagoPileTeam.png License: Public Domain Contributors: Ary29, Davepape, DroEsperanto, Fastfission, Mutter Erde, Wutsje, 2 anonymous edits Image:Atomic cloud over Hiroshima.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Atomic_cloud_over_Hiroshima.jpg License: Public Domain Contributors: Personel aboard Necessary Evil Image:First four nuclear lit bulbs.jpeg Source: http://en.wikipedia.org/w/index.php?title=File:First_four_nuclear_lit_bulbs.jpeg License: Public Domain Contributors: Howcheng, Thuresson, Tungsten, Werewombat, 1 anonymous edits Image:US and USSR nuclear stockpiles.svg Source: http://en.wikipedia.org/w/index.php?title=File:US_and_USSR_nuclear_stockpiles.svg License: Public Domain Contributors: User:Fastfission Image:Pichblende.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Pichblende.jpg License: Creative Commons Attribution-Sharealike 2.5 Contributors: Original uploader was Kgrr at en.wikipedia Image:Citrobacter freundii.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Citrobacter_freundii.jpg License: Public Domain Contributors: Copydays, Georgeryp, Kookaburra, NEON ja, Romary, 1 anonymous edits Image:Yellowcake.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Yellowcake.jpg License: Public Domain Contributors: Mattes, Pfctdayelise, Tungsten Image:Uranium (mined)2.PNG Source: http://en.wikipedia.org/w/index.php?title=File:Uranium_(mined)2.PNG License: Public Domain Contributors: User:Anwar_saadat/bubble_maps_(FAQ) Image:U3O8lattice.jpg Source: http://en.wikipedia.org/w/index.php?title=File:U3O8lattice.jpg License: Public Domain Contributors: Original uploader was Cadmium at en.wikipedia Image:Uranium pourdaix diagram in water.png Source: http://en.wikipedia.org/w/index.php?title=File:Uranium_pourdaix_diagram_in_water.png License: Public Domain Contributors: Cadmium, 1 anonymous edits Image:Uranium pourdiax diagram in carbonate media.png Source: http://en.wikipedia.org/w/index.php?title=File:Uranium_pourdiax_diagram_in_carbonate_media.png License: Public Domain Contributors: Cadmium Image:Uranium fraction diagram with no carbonate.png Source: http://en.wikipedia.org/w/index.php?title=File:Uranium_fraction_diagram_with_no_carbonate.png License: Public Domain Contributors: Cadmium Image:Uranium fraction diagram with carbonate present.png Source: http://en.wikipedia.org/w/index.php?title=File:Uranium_fraction_diagram_with_carbonate_present.png License: Public Domain Contributors: Cadmium File:Uranium-hexafluoride-2D-V2.svg Source: http://en.wikipedia.org/w/index.php?title=File:Uranium-hexafluoride-2D-V2.svg License: unknown Contributors: User:Alexparent, User:Bryan Derksen Image:Gas centrifuge cascade.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Gas_centrifuge_cascade.jpg License: Public Domain Contributors: Fastfission, Pieter Kuiper, 1 anonymous edits
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25
Neptunium
1
Neptunium uranium ← neptunium → plutonium Pm ↑ Np ↓ (Uqt) Periodic Table Extended Periodic Table General Name, symbol, number
neptunium, Np, 93
Element category
actinides
Group, Period, Block
n/a, 7, f
Appearance
silvery metallic Standard atomic weight
−1
(237) g·mol
Electron configuration
4
1
2
[Rn] 5f 6d 7s
Electrons per shell
2, 8, 18, 32, 22, 9, 2 Physical properties
Phase
solid Density (near r.t.)
20.45
Melting point
[1]
−3
g·cm
910 K (637 °C, 1179 °F) Boiling point
4273 K (4000 °C, 7232 °F)
Heat of fusion
3.20 kJ·mol−1
Heat of vaporization
336 kJ·mol−1
Specific heat capacity
(25 °C) 29.46 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
at T(K)
2194
2437
100
1k
Atomic properties Crystal structure Oxidation states
3 forms: orthorhombic, tetragonal and cubic 7, 6, 5, 4, 3 (amphoteric oxide)
10 k
100 k
Neptunium
2 Electronegativity
1.36 (Pauling scale)
Ionization energies
1st: 604.5 kJ/mol
Atomic radius
155 pm
Covalent radius
190±1 pm Miscellaneous [2]
Magnetic ordering
paramagnetic
Electrical resistivity
(22 °C) 1.220 µΩ·m
Thermal conductivity
(300 K) 6.3 W·m−1·K−1
CAS registry number
7439-99-8 Most-stable isotopes
Main article: Isotopes of neptunium iso 235
Np
236
Np
NA syn
syn
half-life 396.1 d
5
1.54×10 y
DM
5.192
231
ε
0.124
235
0.940
236
0.940
236
5.020
232
4.959
233
ε −
α Np
syn
6
2.144×10 y
DP
α
β
237
DE (MeV)
SF & α
Pa U U Pu Pa Pa
References
Neptunium (pronounced /nɛpˈtjuːniəm/) is a chemical element with the symbol Np and atomic number 93. A radioactive metallic element, neptunium is the first transuranic element and belongs to the actinide series. Its most stable isotope, 237Np, is a by-product of nuclear reactors and plutonium production and it can be used as a component in neutron detection equipment. Neptunium is also found in trace amounts in uranium ores due to transmutation reactions.[3]
Characteristics Silvery in appearance, neptunium metal is fairly chemically reactive and is found in at least three allotropes:[3] • α-neptunium, orthorhombic, density 20.45 g/cm³ • β-neptunium (above 280 °C), tetragonal, density (313 °C) 19.36 g/cm³ • γ-neptunium (above 577 °C), cubic, density (600 °C) 18 g/cm³
Neptunium
Compounds This element has four ionic oxidation states while in solution: • • • •
Np3+ (pale purple), analogous to the rare earth ion Pm3+ Np4+ (yellow green) NpO2+ (green blue) NpO22+ (pale pink)
Neptunium forms tri- and tetrahalides such as NpF3, NpF4, NpCl4, NpBr3, NpI3, and oxides of the various compositions such as are found in the uranium-oxygen system, including Np3O8 and NpO2. Neptunium(V) fluoride, NpF5, is volatile like uranium hexafluoride. Neptunium, like other actinides, readily forms a dioxide neptunyl core (NpO2), which readily complexes with carbonate as well as other oxygen moieties (OH−, NO2−, NO3−, and SO42−) to form charged complexes which tend to be readily mobile with low affinities to soil. • NpO2(OH)2− • NpO2(CO3)−
• NpO2(CO3)23− • NpO2(CO3)35−
Uses Precursor in Plutonium-238 Production 237
Np is irradiated with neutrons to create 238Pu, an alpha emitter for radioisotope thermal generators for spacecraft and military applications. 237Np will capture a neutron to form 238 Np and beta decay with a half life of two days to 238Pu.[4]
238
Pu also exists in sizable quantities in spent nuclear fuel but would have to be separated from other isotopes of plutonium.
Weapons applications Neptunium is fissionable, and could theoretically be used as fuel in a fast neutron reactor or a nuclear weapon. In 1992, the U.S. Department of Energy declassified the statement that Np-237 "can be used for a nuclear explosive device".[5] It is not believed that an actual weapon has ever been constructed using neptunium. Calculations show that the critical mass is between 50 and 60 kg.[1] As of 2009, the world production of Np-237 by commercial power reactors was over 1000 critical masses a year, but to extract the isotope from irradiated fuel elements would be a major industrial undertaking. In September 2002, researchers at the University of California Los Alamos National Laboratory created the first known nuclear critical mass using neptunium in combination with enriched uranium, discovering that the critical mass of neptunium is around 60 kg[6] , showing that it "is about as good a bomb material as U-235." US officials in March 2004, planned to move the nation's supply of separated neptunium to a site in Nevada.
3
Neptunium
4
Other 237
Np is used in devices detecting high-energy (MeV) neutrons.[7]
History Neptunium (named for the planet Neptune, the next planet out from Uranus, after which uranium was named) was first discovered by Edwin McMillan and Philip H. Abelson in the year 1940 in Berkeley, California.[8] Initially predicted by Walter Russell's "spiral" organization of the periodic table, it was found at the Berkeley Radiation Laboratory of the University of California, Berkeley where the team produced the neptunium isotope 239Np (2.4 day half-life) by bombarding uranium with slow moving neutrons. It was the first transuranium element produced synthetically and the first actinide series transuranium element discovered.
Occurrence Trace amounts of neptunium are found naturally as decay products from transmutation reactions in uranium ores.[3] Artificial 237Np is produced through the reduction of 237NpF3 with barium or lithium vapor at around 1200 °C and is most often extracted from spent nuclear fuel rods as a by-product in plutonium production. 2 NpF3 + 3 Ba → 2 Np + 3 BaF2 By weight, neptunium-237 discharges are about 5 % as great as plutonium discharges and about 0.05 % of spent nuclear fuel discharges.[9]
Synthesis Chemically, neptunium is prepared by the reduction of NpF3 with barium or lithium vapor at about 1200 °C,[3] however, most Np is produced in nuclear reactions: • When an 235U atom captures a neutron, it is converted to an excited state of 236U. About 81 % of the excited 236U nuclei undergo fission, but the remainder decay to the ground state of 236U by emitting gamma radiation. Further neutron capture creates 237U which has a half-life of 7 days and thus quickly decays to 237Np.
•
237
U is also produced via an (n,2n) reaction with energetic neutrons. • 237Np is the product of alpha decay of 241Am.
238
U. This only happens with very
Heavier isotopes of neptunium decay quickly, and lighter isotopes of neptunium cannot be produced by neutron capture, so chemical separation of neptunium from cooled spent nuclear fuel gives nearly pure 237Np.
Neptunium
Role in nuclear waste Neptunium-237 is the most mobile actinide in the deep geological repository environment.[10] This makes it and its predecessors such as americium-241 candidates of interest for destruction by nuclear transmutation.[11] Neptunium accumulates in commercial household ionization-chamber smoke detectors from decay of the (typically) 0.2 microgram of americium-241 initially present as a source of ionizing radiation. With a half-life of 432 years, the americium-241 in a smoke detector includes about 5 % neptunium after 22 years, and about 10 % after 43 years. After the 432-year americium-241 half-life, a smoke detector's original americium would be almost half neptunium. Due to its long half life neptunium becomes the major contributor of the total radiation in 10000 years. As it is unclear what happens to the containment in that long time span, an extraction of the neptunium would minimize the contamination of the environment if the [12] [13] nuclear waste could be mobilized after several thousand years.
Isotopes 19 neptunium radioisotopes have been characterized, with the most stable being 237Np with a half-life of 2.14 million years, 236Np with a half-life of 154,000 years, and 235Np with a half-life of 396.1 days. All of the remaining radioactive isotopes have half-lives that are less than 4.5 days, and the majority of these have half-lives that are less than 50 minutes. This element also has 4 meta states, with the most stable being 236mNp (t½ 22.5 hours). The isotopes of neptunium range in atomic weight from 225.0339 u (225Np) to 244.068 u (244Np). The primary decay mode before the most stable isotope, 237Np, is electron capture (with a good deal of alpha emission), and the primary mode after is beta emission. The primary decay products before 237Np are element 92 (uranium) isotopes (alpha emission produces element 91, protactinium, however) and the primary products after are element 94 (plutonium) isotopes. 237
Np is fissionable.[6] 237Np eventually decays to form bismuth-209, unlike most other common heavy nuclei which decay to make isotopes of lead. This decay chain is known as the neptunium series.
Literature • Guide to the Elements - Revised Edition, Albert Stwertka, (Oxford University Press; 1998) ISBN 0-19-508083-1 • Lester R. Morss, Norman M. Edelstein, Jean Fuger (Hrsg.): The Chemistry of the Actinide and Transactinide Elements, Springer-Verlag, Dordrecht 2006, ISBN 1-4020-3555-1. • Ida Noddack: "Über das Element 93", in: Angewandte Chemie 1934, 47, 653–655.
5
Neptunium
6
External links • WebElements.com – Neptunium
[14]
(also used as a reference)
• Lab builds world's first neptunium sphere [15], U.S. Department of Energy Research News • NLM Hazardous Substances Databank – Neptunium, Radioactive [16] • Neptunium: Human Health Fact Sheet [17] • C&EN: It's Elemental: The Periodic Table – Neptunium [18] pnb:مینوچپین
References [1] Criticality of a 237Np Sphere (http:/ / typhoon. jaea. go. jp/ icnc2003/ Proceeding/ paper/ 2. 14_107. pdf) [2] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [3] C. R. Hammond (2004). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304857. [4] Lange, R (2008). "Review of recent advances of radioisotope power systems". Energy Conversion and Management 49: 393–401. doi: 10.1016/j.enconman.2007.10.028 (http:/ / dx. doi. org/ 10. 1016/ j. enconman. 2007. 10. 028). [5] "Restricted Data Declassification Decisions from 1946 until Present" (http:/ / www. fas. org/ sgp/ othergov/ doe/ rdd-7. html), accessed Sept 23, 2006 [6] Weiss, P. (October 26, 2002). " Little-studied metal goes critical - Neptunium Nukes? (http:/ / www. findarticles. com/ p/ articles/ mi_m1200/ is_17_162/ ai_94011322)". Science News. . Retrieved 2006-09-29. [7] D. N. Poenaru, Walter Greiner (1997). Experimental techniques in nuclear physics. Walter de Gruyter. p. 236. ISBN 3110144670. [8] Mcmillan, Edwin (1940). "Radioactive Element 93". Physical Review 57: 1185. doi: 10.1103/PhysRev.57.1185.2 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 57. 1185. 2). [9] " Separated Neptunium 237 and Americium (http:/ / www. isis-online. org/ publications/ fmct/ book/ New chapter 5. pdf)" (PDF). . Retrieved 2009-06-06. [10] " Yucca Mountain (http:/ / www. fas. org/ sgp/ othergov/ doe/ lanl/ pubs/ 00818052. pdf)". . Retrieved 2009-06-06. [11] Rodriguez, C (2003). "Deep-Burn: making nuclear waste transmutation practical". Nuclear Engineering and Design 222: 299. doi: 10.1016/S0029-5493(03)00034-7 (http:/ / dx. doi. org/ 10. 1016/ S0029-5493(03)00034-7). [12] Yarris, Lynn (2005-11-29). " Getting the Neptunium out of Nuclear Waste (http:/ / newscenter. lbl. gov/ feature-stories/ 2005/ 11/ 29/ getting-the-neptunium-out-of-nuclear-waste/ )". Berkley laboratory, U.S. Department of Energy. . Retrieved 05-12-2008. [13] " Existing Evidence for the Fate of Neptunium in the Yucca Mountain Repository (http:/ / www. pnl. gov/ main/ publications/ external/ technical_reports/ PNNL-14307. pdf)". Pacific northwest national laboratory, U.S. Department of Energy. January 06-2003. . Retrieved 05-12-2008. [14] [15] [16] [17] [18]
http:/ / www. webelements. com/ neptunium/ http:/ / www. eurekalert. org/ features/ doe/ 2001-08/ danl-lbw060502. php http:/ / toxnet. nlm. nih. gov/ cgi-bin/ sis/ search/ r?dbs+ hsdb:@term+ @na+ @rel+ neptunium,+ radioactive http:/ / www. ead. anl. gov/ pub/ doc/ neptunium. pdf http:/ / pubs. acs. org/ cen/ 80th/ neptunium. html
Article Sources and Contributors
Article Sources and Contributors Neptunium Source: http://en.wikipedia.org/w/index.php?oldid=308828187 Contributors: Addshore, Ahoerstemeier, AlimanRuna, AndonicO, Andres, Andrewa, Arkuat, Avm1, Barticus88, Bayou Banjo, Beetstra, Benbest, Bird, Borislav Dopudja, Breno, Bryan Derksen, Cadmium, Cholmes75, Conversion script, Daggerstab, Darrien, David Latapie, DeadEyeArrow, Deanos, Deconstructhis, Discospinster, DocWatson42, Donarreiskoffer, DragonflySixtyseven, Edgar181, Edward J. Picardy, Egil, El C, Emperorbma, Falcanary, Farseer, Father McKenzie, Femto, Fennec, Fuzzform, Gene Nygaard, Gilliam, Give Peace A Chance, Goldenfool, GraemeL, Grimwald13, Hak-kâ-ngìn, Hallpriest9, HazyM, Hqb, Icairns, Inter, Itinerant1, IvanLanin, JWB, JWBE, Jaraalbe, Jhinman, Jiang, Joanjoc, JoanneB, John, Kazvorpal, Keegan, Kelovy, Kimse, Kuratowski's Ghost, Kwamikagami, LA2, LarryMorseDCOhio, Marc Venot, Materialscientist, Mav, Mboverload, Mdf, Mixwell, Morwen, Neil916, Nergaal, Novangelis, PeepP, Polonium, Poolkris, Pras, PsychoCola, Raeky, Razorflame, RexNL, Reza kalani, Riana, Roberta F., Roentgenium111, Rwendland, Ryukichiro, Santăr, Saperaud, Schneelocke, Sengkang, Sfuerst, Shanedidona, SimonP, Sl, Spiff, Squids and Chips, Stephenb, Stone, Stratocracy, Tagishsimon, Tetracube, Trogdor57, VASANTH S.N., Versus22, Virekleatherwong, Volland, Vsmith, Vuerqex, Warut, Yekrats, Yyy, 146 anonymous edits
Image Sources, Licenses and Contributors image:Np-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Np-TableImage.png License: GNU Free Documentation License Contributors: User:Ahoerstemeier
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7
Plutonium
1
Plutonium neptunium ← plutonium → americium Sm ↑ Pu ↓ (Uqq) Periodic Table Extended Periodic Table General Name, symbol, number
plutonium, Pu, 94
Element category
actinides
Group, Period, Block
n/a, 7, f
Appearance
silvery white
Standard atomic weight
(244) g·mol
Electron configuration
[Rn] 5f 7s
−1
6
Electrons per shell
2
2, 8, 18, 32, 24, 8, 2 Physical properties
Phase
solid Density (near r.t.)
19.816 g·cm−3
Liquid density at m.p.
16.63 g·cm−3
Melting point
912.5 K (639.4 °C, 1182.9 °F) Boiling point
3505 K (3228 °C, 5842 °F)
Heat of fusion
2.82 kJ·mol−1
Heat of vaporization
333.5 kJ·mol−1
Specific heat capacity
(25 °C) 35.5 J·mol−1·K−1 Vapor pressure
P(Pa)
1
10
100
1k
10 k
100 k
at T(K)
1756
1953
2198
2511
2926
3499
Atomic properties
Plutonium
2 Crystal structure
monoclinic
Oxidation states
7, 6, 5, 4, 3 (amphoteric oxide)
Electronegativity
1.28 (Pauling scale)
Ionization energies
1st: 584.7 kJ/mol
Atomic radius
159 pm
Covalent radius
187±1 pm Miscellaneous [1]
Magnetic ordering
paramagnetic
Electrical resistivity
(0 °C) 1.460 µΩ·m
Thermal conductivity
(300 K) 6.74 W·m−1·K−1
Thermal expansion
(25 °C) 46.7 µm·m
Speed of sound (thin rod)
(20 °C) 2260 m/s
Young's modulus
96 GPa
Shear modulus
43 GPa
−1
Poisson ratio
−1
·K
0.21 CAS registry number
7440-07-5 Most-stable isotopes
Main article: Isotopes of plutonium iso 238
Pu
NA syn
half-life 88 y
DM SF α
239
Pu
syn
4
2.41 × 10 y
SF α
240
Pu
241
Pu
syn
syn
6.5 × 103 y
14 y
SF
Pu
syn
3.73 × 105 y
Pu
trace
8.08 × 107 y
5.5 — 5.245 —
— 234
U
— 235
U
—
5.256
236
β−
0.02078
241
SF α
244
—
DP
α
SF 242
DE (MeV)
α SF
U Am
—
—
—
—
4.984
238
4.666
240
—
U U
—
References
Plutonium (pronounced /pluːˈtoʊniəm/, symbol Pu, atomic number—or element—94) is a rare transuranic radioactive element. It is an actinide metal of silvery-white appearance that tarnishes when exposed to air, forming a dull coating when oxidized. The element
Plutonium normally exhibits six allotropes and four oxidation states. It reacts with carbon, halogens, nitrogen and silicon. When exposed to moist air, it forms oxides and hydrides that expand the sample up to 70% in volume, which in turn flake off as a powder that can spontaneously ignite. It is also a radioactive poison that accumulates in bone marrow. These and other properties make the handling of plutonium dangerous, although its overall toxicity is sometimes overstated. • Plutonium-238 has a half-life of 88 years and emits alpha particles. It is a heat source in radioisotope thermoelectric generators, which are used to power some spacecraft. Pu-238 is formed in several ways, but pure Pu-238 for RTGs is produced by neutron capture on neptunium-237. • Plutonium-239 is the most important isotope of plutonium, with a half-life of 24,100 years. Pu-239 and Pu-241 are fissile, meaning that the nuclei of its atoms can break apart by being bombarded by slow moving thermal neutrons, releasing energy, gamma radiation and more neutrons. It can therefore sustain a nuclear chain reaction, leading to applications in nuclear weapons and nuclear reactors. Pu-239 is synthesized by irradiating uranium-238 with neutrons in a nuclear reactor, then recovered via nuclear reprocessing of the fuel. Further neutron capture produces successively heavier isotopes. • Plutonium-240 has a high rate of spontaneous fission, raising the background neutron radiation of plutonium containing it. Plutonium is graded by proportion of Pu-240: weapons grade (< 7%), fuel grade (7–19%) and reactor grade (> 19%). Lower grades are less suited for nuclear weapons and thermal reactors but can fuel fast reactors. Pu-240 is not fissile, but is fertile material like U-238. • Plutonium-241 is fissile, but also beta decays with a halflife of 14 years to americium-241. • Plutonium-242 is not fissile, not very fertile (requiring 3 more neutron captures to become fissile), has a low neutron capture cross section, and a longer halflife than any of the lighter isotopes. • Plutonium-244 is the most stable isotope of plutonium, with a half-life of about 80 million years, long enough to be found in trace quantities in nature. Element 94 was first synthesized in 1940 by a team led by Glenn T. Seaborg and Edwin McMillan at a University of California, Berkeley laboratory by bombarding uranium-238 with deuterons. McMillan named the new element after Pluto, and Seaborg suggested the symbol Pu as a joke. Trace amounts of plutonium were subsequently discovered in nature. Discovery of plutonium became a classified part of the Manhattan Project to develop an atomic bomb during World War II. The first nuclear test, "Trinity" (July 1945), and the second atomic bomb used to destroy a city (Nagasaki, Japan, in August 1945), "Fat Man", both had cores of Pu-239. Human radiation experiments studying plutonium were conducted without informed consent, and a number of criticality accidents, some lethal, occurred during and after the war. Disposal of plutonium waste from nuclear power plants and dismantled nuclear weapons built during the Cold War is a major nuclear-proliferation, health, and environmental concern. Other sources of plutonium in the environment are fallout from numerous above-ground nuclear tests (now banned) and several nuclear accidents.
3
Plutonium
4
Characteristics Physical Plutonium, like most metals, has a bright silvery appearance at first, much like nickel, but it oxidizes very quickly to a dull gray, although yellow and olive green are also reported.[2] [3] At room temperature plutonium is in its α form (alpha). This, the most common structural form of the element (allotrope), is about as hard and brittle as grey cast iron unless it is alloyed with other metals to make it soft and ductile.[2] Unlike most metals, it is not a good conductor of heat or electricity.[2] It has a low melting point (640 °C) and an unusually high boiling point (3,327 °C).[2] Alpha particle emission, which is the release of high-energy helium nuclei, is the most common form of radiation given off by plutonium.[4] Heat given off by the release of and deceleration of these alpha particles make a mass of plutonium the size of a softball warm to the touch while a somewhat larger mass can boil a liter of water in a few minutes, although this varies with isotopic composition.[5] [6] Resistivity is a measure of how strongly a material opposes the flow of electric current. The resistivity of plutonium at room temperature is very high for a metal, and it gets even higher with lower temperatures, which is unusual for metals.[7] This trend continues down to 100 K, below which resistivity rapidly decreases for fresh samples.[7] Resistivity then begins to increase with time at around 20 K due to radiation damage, with the rate dictated by the isotopic composition of the sample.[7] Due to self-irradiation, a sample of plutonium fatigues throughout its crystal structure, meaning the ordered arrangement of its atoms becomes disrupted by radiation with time.[8] However, self-irradiation can also lead to annealing which counteracts some of the fatigue effects as temperature increases above 100 K.[9]
Allotropes Plutonium normally has six allotropes and forms a seventh (zeta, ζ) at high temperature within a limited pressure range.[10] These allotropes, which are different structural modifications or forms of an element, have very similar internal energies but significantly varying densities and crystal structures. This makes plutonium very sensitive to changes in temperature, pressure, or chemistry, and allows for dramatic volume changes following phase transitions from one allotropic form to another.[8] Unlike most materials, plutonium increases in density when it melts, by 2.5%, but the liquid metal exhibits a linear decrease in density with temperature.[7] Densities of the different allotropes vary from 16.00 g/cm3 to 19.86 g/cm3.[11] The presence of these many allotropes makes machining plutonium very difficult, as it changes state
Plutonium has six allotropes at ambient pressure: alpha (α), beta (β), gamma (γ), delta (δ), [10] delta prime (δ′), & epsilon (ε)
Plutonium
5
very readily. For example, the α form exists at room temperature in unalloyed plutonium. It has machining characteristics similar to cast iron but changes to the plastic and malleable β form (beta) at slightly higher temperatures.[12] The reasons for the complicated phase diagram are not entirely understood. The α form has a low-symmetry monoclinic structure, [10] hence its brittleness, strength, compressibility, and poor conductivity. Plutonium in the δ form (delta) normally exists in the 310 °C to 452 °C range but is stable at room temperature when alloyed with a small percentage of gallium, aluminium, or cerium, enhancing workability and allowing it to be welded.[12] The delta form has more typical metallic character, and is roughly as strong and malleable as aluminium.[10]
Nuclear fission Plutonium is a radioactive actinide metal whose isotope, plutonium-239 (Pu-239), is one of the three primary fissile isotopes[13] (uranium-233 and uranium-235 are the other two).[14] Pu-241 is also highly fissile. To be considered fissile, an isotope's atomic nucleus must be able to break apart or fission when struck by a slow moving neutron, and to release enough additional neutrons in the process to sustain the nuclear chain reaction by splitting further nuclei. Pu-239 has a multiplication factor (k) larger than one, which means
Weapons-grade electrorefined plutonium
that if the metal is present in sufficient mass and with an appropriate geometry (e.g., a compressed sphere), it can form a critical mass.[15] During fission, a fraction of the binding energy, which holds a nucleus together, is released as a large amount of thermal, electromagnetic and kinetic energy; a kilogram of Pu-239 can produce an explosion equivalent to 20,000 tons of TNT.[5] It is this energy that makes Pu-239 useful in nuclear weapons and reactors. The presence of the isotope plutonium-240 (Pu-240) in a sample limits
its nuclear bomb potential, as Pu-240 has a relatively high spontaneous fission rate (~440 fissions per second per gram—over 1,000 neutrons per second per gram[16] ), raising the background neutron levels and thus increasing the risk of predetonation.[17] Plutonium is identified as either weapon grade, fuel grade, or power reactor grade based on the percentage of Pu-240 that it contains. Weapon grade plutonium contains less than 7% Pu-240. Fuel grade plutonium contains from 7 to less than 19%, and power reactor grade contains 19% or more Pu-240.[18] The isotope plutonium-238 (Pu-238) is not fissile but can undergo nuclear fission easily with fast neutrons as well as alpha decay.[5]
Isotopes and synthesis Twenty radioactive isotopes of plutonium have been characterized.[4] The longest-lived are Pu-244, with a half-life of 80.8 million years, Pu-242, with a half-life of 373,300 years, and Pu-239, with a half-life of 24,110 years.[4] All of the remaining radioactive isotopes have half-lives that are less than 7,000 years.[4] This element also has eight metastable states, though none are stable and all have half-lives less than one second.[4] The isotopes of plutonium range in mass number from 228 to 247.[4] The primary decay modes of isotopes with mass numbers lower than the most stable isotope, Pu-244, are spontaneous fission and α emission, mostly forming uranium (92 protons) and neptunium
Plutonium
6
(93 protons) isotopes as decay products (neglecting the wide range of daughter nuclei created by fission processes).[4] The primary decay mode for isotopes with mass numbers higher than Pu-244 is β emission, mostly forming americium (95 protons) isotopes as decay products.[4] Pu-241 is the parent isotope of the neptunium decay series, decaying to [5] americium-241 via β or electron emission. Pu-238 and Pu-239 are the most-widely synthesized isotopes.[5] Pu-239 is synthesized via the following reaction using uranium (U) and neutrons (n) via beta decay (β−) with neptunium (Np) as an intermediate:[19]
In other words, neutrons from the fission of U-235 are captured by U-238 nuclei to form U-239; a beta decay converts a neutron into a proton to form Np-239 (half-life 2.36 days) and another beta decay forms Pu-239.[20] Workers on the Tube Alloys project had predicted this reaction theoretically in 1940. Pu-238 is synthesized by bombarding U-238 with deuterons (D, the nuclei of heavy hydrogen) in the following reaction:[21]
In this equation, a deuteron hitting U-238 produces two neutrons and Np-238. The Np-238 spontaneously decays by emitting negative beta particles to form Pu-238.
Compounds and chemistry At
room
temperature,
pure
plutonium is silvery in color but gains a tarnish when oxidized.[5] The element displays four common ionic oxidation states in aqueous solution and one rare one:[11] • Pu(III), as Pu3+ (blue lavender) • Pu(IV), as Pu4+ (yellow brown) • Pu(V), as PuO2+ (pink?)[22]
• Pu(VI), as PuO22+ (pink orange) • Pu(VII), as PuO53− (green)–the heptavalent ion is rare
The color shown by plutonium solutions depends on both the Various oxidation states of Pu in solution oxidation state and the nature of [23] the acid anion. It is the acid anion that influences the degree of complexing—how atoms connect to a central atom—of the plutonium species. Metallic plutonium is produced by reacting plutonium(IV) fluoride with barium, calcium or lithium at 1200 °C.[24] It is attacked by acids, oxygen, and steam but not by alkalis and dissolves easily in concentrated hydrochloric, hydroiodic and perchloric acids.[25] Molten metal must be kept in a vacuum or an inert atmosphere to avoid reaction with air.[12] At 135 °C the metal will ignite in air and will explode if placed in carbon tetrachloride.[26]
Plutonium
7 Plutonium is a reactive metal. In moist air or moist argon, the metal oxidizes rapidly, producing a mixture of oxides and hydrides.[2] If the metal is exposed long enough to a limited amount of water vapor, a powdery surface coating of PuO2 is formed.[2] Also formed is plutonium hydride but an excess of water vapor forms only PuO2.[25]
Plutonium pyrophoricity can cause it to look like a glowing ember under certain conditions.
With this coating, the metal is pyrophoric, meaning it can ignite spontaneously, so plutonium metal is usually handled in an inert, dry atmosphere of nitrogen or argon.[2] Oxygen retards the effects of moisture and
acts as a passivating agent.[2] Plutonium reacts readily with oxygen, forming PuO and PuO2 as well as intermediate oxides;[11] plutonium oxide fills 40% more volume than plutonium metal.[26] It reacts with the halogens, giving rise to compounds such as PuX3 where X can be F, Cl, Br or I; PuF4 is also seen.[11] The following oxyhalides are observed: PuOCl, PuOBr and PuOI.[11] It will react with carbon to form PuC, nitrogen to form PuN and silicon to form PuSi2.[11] Crucibles used to contain plutonium need to be able to withstand its strongly reducing properties.[12] Refractory metals such as tantalum and tungsten along with the more stable oxides, borides, carbides, nitrides and silicides can tolerate this.[12] Melting in an electric arc furnace can be used to produce small ingots of the metal without the need for a crucible.[12] Plutonium can form alloys and intermediate compounds with most other metals. Exceptions include lithium, sodium, potassium, and rubidium of the alkali metals; and magnesium, calcium, strontium, and barium of the alkaline earth metals; and europium and ytterbium of the rare earth metals.[25] Partial exceptions include the refractory metals chromium, molybdenum, niobium, tantalum, and tungsten, which are soluble in liquid plutonium, but insoluble or only slightly soluble in solid plutonium.[25]
Occurrence Trace amounts of two plutonium isotopes (Pu-239 and Pu-244) can be found in nature. Small traces of Pu-239, a few parts per trillion, and its decay products are naturally found in some concentrated ores of uranium,[27] such as the natural nuclear fission reactor in Oklo, Gabon.[28] The ratio of Pu-239 to U at the Cigar Lake Mine uranium deposit ranges from 2.4 × 10−12 to 44 × 10−12.[29] Even smaller amounts of primordial Pu-244 occur naturally due to its relatively long half-life of about 80 million years.[30] Minute traces are found in the human body due to the 550 above-ground nuclear tests which have been performed and several major nuclear accidents.[26] Most atmospheric nuclear testing was stopped in 1963 by the Limited Test Ban Treaty but France continued to test into the 1980s and several other nations also conducted tests after 1963. Because it is specifically manufactured and is the result of radioactive decay of uranium ores, Pu-239 is the most abundant isotope of plutonium.[26]
Plutonium
8
History Discovery Enrico Fermi and a team of scientists at the University of Rome reported that they had discovered element 94 in 1934.[31] Fermi called the element hesperium and mentioned it in his Nobel Lecture in 1938.[32] However, the sample was actually a mixture of barium, krypton, and other elements, but this was not known at the time because nuclear fission had not been discovered yet.[33] Plutonium (specifically, Pu-238) was first produced and isolated on December 14, 1940, and chemically identified on February 23, 1941, by Dr. Glenn T. Seaborg, Edwin M. McMillan, J. W. Kennedy, Z. M. Tatom, and A. C. Wahl by deuteron bombardment of uranium in the 60-inch (150 cm) cyclotron at the University of California, Berkeley.[34] In the 1940 experiment, neptunium-238 was created directly by the bombardment but decayed by beta emission two days later, which indicated the formation of element 94.[26] A paper documenting the discovery was prepared by the team and sent to the journal Physical Review in March 1941.[26] The paper was withdrawn before publication after the discovery that an isotope of the new element (Pu-239) could undergo nuclear fission in a way that might be useful in an atomic bomb. Publication was delayed until a year after the end of World War II due to security concerns.[13]
Glenn T. Seaborg and his team at Berkeley were the first to produce plutonium.
Edwin McMillan had recently named the first transuranium element after the planet Neptune and suggested that element 94, being the next element in the series, be named for what was then considered the next planet, Pluto.[5] [35] Seaborg originally considered the name "plutium", but later thought that it did not sound as good as "plutonium."[36] He chose the letters "Pu" as a joke, which passed without notice into the periodic table.[37] [38] [39]
Alternate names considered by Seaborg and others were "ultimium" or "extremium" because of the now-discredited belief that they had found the last possible element on the periodic table.[40]
Early research The basic chemistry of plutonium was found to resemble uranium after a few months of initial study.[26] Early research was continued at the secret Metallurgical Laboratory of the University of Chicago. On August 18, 1942, a trace quantity of this element was isolated and measured for the first time. About 50 micrograms of plutonium-239 combined with uranium and fission products was produced and only about 1 microgram was isolated.[27] This procedure enabled chemists to determine the new element's atomic weight.[41] [42] In November 1943 some plutonium trifluoride was reduced to create the first sample of plutonium metal: a few micrograms of metallic beads.[27] Enough plutonium was produced
Plutonium
9
to make it the first synthetically made element to be visible with the unaided eye.[43] The nuclear properties of plutonium-239 were also studied; researchers found that when it is hit by a neutron it breaks apart (fissions) by releasing more neutrons and energy. These neutrons can hit other atoms of plutonium-239 and so on in an exponentially fast chain reaction. This can result in an explosion large enough to destroy a city if enough of the isotope is concentrated to form a critical mass.[26]
Production during the Manhattan Project During World War II the U.S. government established the Manhattan Project, which was tasked with developing an atomic bomb. The three primary research and production sites of the project were the plutonium production facility at what is now the Hanford Site, the uranium enrichment facilities at Oak Ridge, Tennessee, and the weapons research and design laboratory, now known as Los Alamos National Laboratory.[44] The first production reactor that made plutonium-239 was the X-10 Graphite Reactor. It went online in 1943 and was built at a facility in Oak Ridge that later became the Oak Ridge National Laboratory.[26] [45] [46] [47]
On April 5, 1944, Emilio Segrè at Los Alamos received the first sample of reactor-produced plutonium from Oak Ridge.[48] Within ten days, he discovered that The Hanford B Reactor face under reactor-bred plutonium had a higher concentration of construction—the first plutonium-production reactor the isotope Pu-240 than cyclotron-produced plutonium. Pu-240 has a high spontaneous fission rate, raising the overall background neutron level of the plutonium sample. The original gun-type plutonium weapon, code-named "Thin Man", had to be abandoned as a result—the increased number of spontaneous neutrons meant that nuclear pre-detonation (a fizzle) would be likely. The entire plutonium weapon design effort at Los Alamos was soon changed to the more complicated implosion device, code-named "Fat Man." With an implosion weapon, a solid sphere of plutonium is compressed to a high density with explosive lenses—a technically more daunting task than the simple gun-type design, but necessary in order to use plutonium for weapons purposes. (Enriched uranium, by contrast, can be used with either method.)[48] Construction of the Hanford B Reactor, the first industrial-sized nuclear reactor for the purposes of material production, was completed in March 1945.[49] B Reactor produced the fissile material for the plutonium weapons used during World War II.[50] [51] [52] [53] [54]
B, D and F were the initial reactors built at Hanford, and six additional plutonium-producing reactors were built later at the site.[49] In 2004, a safe was discovered during excavations of a burial trench at the Hanford nuclear site. Inside the safe were various items, including a large glass bottle containing a whitish
Plutonium
10
slurry which was subsequently identified as the oldest sample of weapons-grade plutonium known to exist. Isotope analysis by Pacific Northwest National Laboratory indicated that the plutonium in the bottle was manufactured in the X-10 reactor at Oak Ridge during 1944. [55] [56] [57]
Trinity and Fat Man atomic bombs The first atomic bomb test, codenamed "Trinity" and detonated on July 16, 1945, near Alamogordo, New Mexico, used plutonium as its fissile material.[27] The implosion design of "the Gadget", as the Trinity device was code-named, used conventional explosive lenses to compress a sphere of plutonium into a supercritical mass, which was simultaneously showered with neutrons from an initiator made of polonium and beryllium (neutron source: (α, n) reaction).[26] Together, these ensured a runaway chain reaction and explosion. The overall weapon weighed over 4 tonnes, although it used just 6.2 kg of plutonium in its core.[58] About 20% of the plutonium used in the Trinity weapon underwent fission, resulting in an explosion with an energy equivalent to approximately 20,000 tons of TNT.[59] [60]
Because of the presence of Pu-240 in reactor-bred plutonium, the implosion design was developed for the "Fat Man" and Trinity" weapons
[61] [61]
An identical design was used in the "Fat Man" atomic bomb dropped on Nagasaki, Japan, on August 9, 1945, killing 70,000 people and wounding another 100,000.[26] The "Little Boy" bomb dropped on Hiroshima three days earlier used uranium-235, not plutonium. Japan capitulated on August 15 to General Douglas MacArthur, effectively ending the war. Only after the announcement of the first atomic bombs was the existence of plutonium made public.
Cold War use and waste Large stockpiles of weapons-grade plutonium were built up by both the Soviet Union and the United States during the Cold War. The U.S. reactors at Hanford and the Savannah River Site in South Carolina produced 103 tonnes,[62] and an estimated 170 tonnes of military-grade plutonium was produced in Russia.[63] [64] Each year about 20 tonnes of the element is still produced as a by-product of the nuclear power industry.[11] As much as 1000 tonnes of plutonium may be in storage with more than 200 tonnes of that either inside or extracted from nuclear weapons.[26] SIPRI estimated the world plutonium stockpile in 2007 as about 500 tons, divided equally between weapon and civilian stocks, but all weapon-usable..[65]
Plutonium
Since the end of the Cold War, these stockpiles have become a focus of nuclear proliferation concerns. In the U.S., some plutonium extracted from dismantled nuclear weapons is melted to form glass logs of plutonium oxide that weigh two tonnes.[26] The glass is made of borosilicates mixed with cadmium and gadolinium.[66] These logs are planned to be encased in stainless steel and stored as much as 4 km underground Proposed waste storage tunnel design in bore holes that will be back-filled with concrete.[26] for the Yucca Mountain nuclear waste As of 2008, the only facility in the U.S. that is scheduled repository to store plutonium in this way is the Yucca Mountain nuclear waste repository, which is about 100 miles (160 km) north-east of Las Vegas, Nevada.[67] Local and state opposition to this plan has delayed efforts to store nuclear waste at Yucca Mountain.
Medical experimentation During and after the end of World War II, scientists working on the Manhattan Project and other nuclear weapons research projects conducted studies of the effects of plutonium on laboratory animals and human subjects.[68] Animal studies found that a few milligrams of plutonium per kilogram of tissue is a lethal dose.[69] In the case of human subjects, this involved injecting solutions containing (typically) five micrograms of plutonium into hospital patients thought to be either terminally ill, or to have a life expectancy of less than ten years either due to age or chronic disease condition.[68] This was reduced to one microgram in July 1945 after animal studies found that the way plutonium distributed itself in bones was more dangerous than radium.[69] Eighteen human test subjects were injected with plutonium without informed consent.[68] The tests were used to create diagnostic tools to determine the uptake of plutonium in the body in order to develop safety standards for working with plutonium.[68] The episode is now considered to be a serious breach of medical ethics and of the Hippocratic Oath. More sympathetic commentators have noted that while it was definitely a breach in trust and ethics, "the effects of the plutonium injections were not as damaging to the subjects as the early news stories painted, nor were they so inconsequential as many scientists, then and now, believe."[70]
11
Plutonium
Applications Explosives The isotope Pu-239 is a key fissile component in nuclear weapons, due to its ease of fission and availability. Encasing the bomb's sphere of plutonium in a tamper (an optional layer of dense material) decreases the amount of plutonium needed to reach critical mass by reflecting escaping neutrons back into the plutonium core. This reduces the amount of plutonium needed to reach criticality from 16 kg to 10 kg, which is a sphere with a diameter of about 10 centimetres (4 in).[71] This critical mass is about a third of that for U-235.[5] The "Fat Man"-type plutonium bombs produced during the The atomic bomb dropped on Manhattan Project used explosive compression of plutonium Nagasaki, Japan in 1945 had a to obtain significantly higher densities than normal, plutonium core. combined with a central neutron source to begin the reaction and increase efficiency. Thus only 6.2 kg of plutonium was needed for an explosive yield equivalent to 20 kilotons of TNT.[59] [72] (See also Nuclear weapon design.) Hypothetically, as little as 4 kg of plutonium—and maybe even less—could be used to make a single atomic bomb using very sophisticated assembly designs.[72]
Use of nuclear waste PUREX (Plutonium–URanium EXtraction) reprocesses spent nuclear fuel to extract uranium and plutonium to form a mixed oxide (MOX) fuel for reuse in nuclear reactors. Weapons grade plutonium can be added to the fuel mix. MOX fuel is used in light water reactors and consists of 60 kg of plutonium per tonne of fuel; after four years, three-quarters of the plutonium is burned (turned into other elements).[26] Breeder reactors are specifically designed to create more fissionable material than they consume. MOX fuel has been in use since the 1980s and is widely used in Europe.[73] In September 2000, the United States and the Russian Federation signed a Plutonium Management and Disposition Agreement by which each agreed to dispose of 34 tonnes of weapon grade plutonium.[74] The U.S. Department of Energy plans to dispose of 34 tonnes of weapon grade plutonium in the United States before the end of 2019 by converting the plutonium to a MOX fuel to be used in commercial nuclear power reactors.[74] Efficiencies are also attained through reprocessing: a fuel rod is reprocessed after three years of use to remove waste products, which by then account for 3% of the total weight of the rods.[26] Any uranium or plutonium isotopes produced during those three years are left and the rod goes back into production.[75] However, the presence of up to 1% gallium per mass in weapon grade plutonium has the potential to interfere with long-term operation of a light water reactor.[76] 241
Am has recently been suggested for use as a denaturing agent in plutonium reactor fuel rods to render the fuel unusable for conversion to nuclear weapons.[77]
12
Plutonium
13
Power and heat source The isotope plutonium-238 (Pu-238) has a half-life of 87.5 years. It emits a large amount of thermal energy with low levels of both gamma rays/particles and spontaneous neutron rays/particles.[78] Being an alpha emitter, it combines high energy radiation with low penetration and thereby requires minimal shielding. A sheet of paper can be used to shield against the alpha particles emitted by Pu-238 while one kilogram of the isotope can generate 22 million kilowatt-hours of heat energy.[5] [78]
A glowing pellet of
238
PuO2
These characteristics make it well-suited for electrical power generation for devices which must function without direct maintenance for timescales approximating a human lifetime. It is therefore used in radioisotope thermoelectric generators and radioisotope heater units such as those in the Cassini, Voyager and New Horizons space probes. Earlier versions of the same technology powered the ALSEP and EASEP systems including seismic experiments on the Apollo 14 Moon mission.[26] Plutonium-238 has also been used successfully to power artificial heart pacemakers, to reduce the risk of repeated surgery.[79] [80] It has been largely replaced by lithium-based primary cells, but as of 2003 there were somewhere between 50 and 100 plutonium-powered pacemakers still implanted and functioning in living patients.[81] Pu-238 was studied as way to provide supplemental heat to scuba diving.[82] Plutonium-238 mixed with beryllium is used to generate neutrons for research purposes.[26]
Precautions Toxicity Isotopes and compounds of plutonium are toxic to highly toxic due to their radioactivity. Contamination by plutonium oxide (spontaneously oxidized plutonium) has resulted from a number of military nuclear accidents where nuclear weapons have burned.[83] However, based on chemical toxicity alone, the element is less dangerous than arsenic or cyanide and about the same as caffeine.[84] [85] Plutonium is more dangerous when inhaled than when ingested. The risk of lung cancer increases once the total dose equivalent of inhaled radiation exceeds 400 mSv.[86] The U.S. Department of Energy estimates that the lifetime cancer risk for inhaling 5,000 plutonium particles, each about 3 microns wide, to be 1% over the background U.S. average.[87] It is not absorbed into the body efficiently when ingested; only 0.04% of plutonium oxide is absorbed after ingestion.[26] When plutonium is absorbed into the body, it is excreted very slowly, with a biological half-life of 200 years.[88] Plutonium has a metallic taste.[89] The alpha radiation plutonium emits does not penetrate the skin but can irradiate internal organs when plutonium is inhaled or ingested.[26] Particularly at risk are the skeleton, where it is likely to be absorbed by the bone surface, and the liver, where it collects and becomes concentrated.[25] Considerably larger amounts may cause acute radiation poisoning and death if ingested or inhaled; however, no human is known to have died because of inhaling or ingesting plutonium, and many people have measurable amounts of plutonium in their bodies.[85]
Plutonium
Criticality potential Toxicity issues aside, care must be taken to avoid the accumulation of amounts of plutonium which approach critical mass, particularly because plutonium's critical mass is only a third of that of uranium-235.[5] A critical mass of plutonium emits lethal amounts of neutrons and gamma rays.[90] Plutonium in solution is more likely to form a critical mass than the solid form due to moderation by the hydrogen in water.[11] Criticality accidents have occurred in the past, some of A simulated sphere of plutonium them with lethal consequences. Careless handling of surrounded by neutron-reflecting tungsten carbide bricks around a 6.2 kg plutonium tungsten carbide blocks in a re-enactment of Harry Daghlian's 1945 sphere resulted in a fatal dose of radiation at Los experiment Alamos on August 21, 1945, when scientist Harry K. Daghlian, Jr. received a dose estimated to be 5.1 Sievert (510 rems) and died 28 days later.[91] Nine months later, another Los Alamos scientist, Louis Slotin, died from a similar accident involving a beryllium reflector and the same plutonium core (the so-called "demon core") that had previously claimed the life of Daghlian.[92] These incidents were fictionalized in the 1989 film Fat Man and Little Boy. In December 1958, during a process of purifying plutonium at Los Alamos, a critical mass was formed in a mixing vessel, which resulted in the death of a crane operator.[93] Other nuclear accidents have occurred in the Soviet Union, Japan, and many other countries.[93]
Flammability Metallic plutonium is a fire hazard, especially if the material is finely divided.[94] In a moist environment, plutonium forms hydrides on its surface, which are pyrophoric and may ignite in air at room temperature.[94] Plutonium expands up to 70% in volume as it oxidizes and thus may break its container.[94] The radioactivity of the burning material is an additional hazard. Magnesium oxide sand is probably the most effective material for extinguishing a plutonium fire.[94] It cools the burning material, acting as a heat sink, and also blocks off oxygen. Special precautions are necessary to store or handle plutonium in any form; generally a dry inert gas atmosphere is required.[95] [96] [97] [98]
14
Plutonium
See also • Nuclear engineering • Nuclear fuel cycle • Nuclear physics
References [1] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/ elementmagn. pdf), in Handbook of Chemistry and Physics 81th edition, CRC press. [2] NIH contributors. " Plutonium, Radioactive (http:/ / webwiser. nlm. nih. gov/ getSubstanceData. do;jsessionid=89B673C34252C77B4C276F2B2D0E4260?substanceID=419& displaySubstanceName=Plutonium, Radioactive& UNNAID=& STCCID=& selectedDataMenuItemID=44)". Wireless Information System for Emergency Responders (WISER). Bethesda (MD): U.S. National Library of Medicine, National Institutes of Health. . Retrieved 2008-11-23. (public domain text) [3] ARQ staff (2008). " Nitric acid processing (http:/ / arq. lanl. gov/ source/ orgs/ nmt/ nmtdo/ AQarchive/ 3rdQuarter08/ page3. shtml)". Actinide Research Quarterly (Los Alamos (NM): Los Alamos National Laboratory) (3rd quarter). . Retrieved 2009-02-15. "While plutonium dioxide is normally olive green, samples can be various colors. It is generally believed that the color is a function of chemical purity, stoichiometry, particle size, and method of preparation, although the color resulting from a given preparation method is not always reproducible.". [4] NNDC contributors (2008). " Chart of Nuclides (http:/ / www. nndc. bnl. gov/ chart/ )". in Alejandro A. Sonzogni (Database Manager). Upton (NY): National Nuclear Data Center, Brookhaven National Laboratory. . Retrieved 2008-09-13. [5] Heiserman 1992 [6] Rhodes, Richard (1986). The Making of the Atomic Bomb. New York: Simon & Schuster. pp. 659–660. ISBN 0-671-65719-4. Leona Marshall: "When you hold a lump of it in your hand, it feels warm, like a live rabbit" [7] Miner 1968, p. 544 [8] Hecker, Siegfried S. (2000). " Plutonium and its alloys: from atoms to microstructure (http:/ / www. fas. org/ sgp/ othergov/ doe/ lanl/ pubs/ 00818035. pdf)" (PDF). Los Alamos Science 26: 290–335. . Retrieved 2009-02-15. [9] Hecker, Siegfried S.; Martz, Joseph C. (2000). " Aging of Plutonium and Its Alloys (http:/ / library. lanl. gov/ cgi-bin/ getfile?00818029. pdf)" (PDF). Los Alamos Science (Los Alamos, New Mexico: Los Alamos National Laboratory) (26): 242. . Retrieved 2009-02-15. [10] Baker, Richard D.; Hecker, Siegfried S.; Harbur, Delbert R. (1983). " Plutonium: A Wartime Nightmare but a Metallurgist's Dream (http:/ / library. lanl. gov/ cgi-bin/ getfile?07-16. pdf)". Los Alamos Science (Los Alamos National Laboratory): 148, 150–151. . Retrieved 2009-02-15. [11] CRC 2006, p. 4–27 [12] Miner 1968, p. 542 [13] Stwertka 1998 [14] EPA contributors (2008). " Fissile Material (http:/ / www. epa. gov/ rpdweb00/ glossary/ termdef. html#f)". Radiation Glossary. United States Environmental Protection Agency. . Retrieved 2008-11-23. [15] Asimov, Isaac (1988). "Nuclear Reactors". Understanding Physics. Barnes & Noble Publishing. p. 905. ISBN 0880292512. [16] Samuel Glasstone and Leslie M. Redman, An Introduction to Nuclear Weapons (http:/ / www. doeal. gov/ opa/ docs/ RR00171. pdf) (Atomic Energy Commission Division of Military Applications Report WASH-1038, June 1972), p. 12. [17] Gosling, F.G. (1999). The Manhattan Project: Making the Atomic Bomb (http:/ / www. cfo. doe. gov/ me70/ manhattan/ publications/ DE99001330. pdf). Oak Ridge (TN): United States Department of Energy. p. 40. DOE/MA-0001-01/99. . Retrieved 2009-02-15. [18] DOE contributors (1996). Plutonium: The First 50 Years (http:/ / www. doeal. gov/ SWEIS/ DOEDocuments/ 004 DOE-DP-0137 Plutonium 50 Years. pdf). U.S. Department of Energy. DOE/DP-1037. . (public domain text) [19] Kennedy, J. W.; Seaborg, G. T.; Segrè, E.; Wahl, A. C. (1946). "Properties of Element 94". Physical Review 70 (7–8): 555–556. doi: 10.1103/PhysRev.70.555 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 70. 555). [20] Greenwood 1997, p. 1259 [21] Seaborg, Glenn T.; McMillan, E.; Kennedy, J. W.; Wahl, A. C. (1946). "Radioactive Element 94 from Deuterons on Uranium". Physical Review 69 (7–8): 366–367. doi: 10.1103/PhysRev.69.367 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 69. 367).
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Plutonium [22] The PuO2+ ion is unstable in solution and will disproportionate into Pu4+ and PuO22+; the Pu4+ will then oxidize the remaining PuO2+ to PuO22+, being reduced in turn to Pu3+. Thus, aqueous solutions of plutonium tend over time towards a mixture of Pu3+ and PuO22+.
Crooks, William J. (2002). " Nuclear Criticality Safety Engineering Training Module 10 – Criticality Safety in Material Processing Operations, Part 1 (http:/ / ncsp. llnl. gov/ ncset/ Module10. pdf)" (PDF). . Retrieved 2006-02-15. [23] Matlack, George (2002). A Plutonium Primer: An Introduction to Plutonium Chemistry and its Radioactivity. Los Alamos National Laboratory. LA-UR-02-6594. [24] Eagleson, Mary (1994). Concise Encyclopedia Chemistry. Walter de Gruyter. p. 840. ISBN 9783110114515. [25] Miner 1968, p. 545 [26] Emsley 2001 [27] Miner 1968, p. 541 [28] DOE contributors (2004). " Oklo: Natural Nuclear Reactors (http:/ / www. ocrwm. doe. gov/ factsheets/ doeymp0010. shtml)". U.S. Department of Energy, Office of Civilian Radioactive Waste Management. . Retrieved 2008-11-16. [29] Curtis, David; Fabryka-Martin, June; Paul, Dixon; Cramer, Jan (1999). "Nature's uncommon elements: plutonium and technetium". Geochimica et Cosmochimica Acta 63 (2): 275–285. doi: 10.1016/S0016-7037(98)00282-8 (http:/ / dx. doi. org/ 10. 1016/ S0016-7037(98)00282-8). [30] Hoffman, D. C.; Lawrence, F. O.; Mewherter, J. L.; and Rourke, F. M. (1971). "Detection of Plutonium-244 in Nature". Nature 234: 132–134. doi: 10.1038/234132a0 (http:/ / dx. doi. org/ 10. 1038/ 234132a0). Nr. 34. [31] Holden, Norman E. (2001). " A Short History of Nuclear Data and Its Evaluation (http:/ / www. nndc. bnl. gov/ content/ evaluation. html)". 51st Meeting of the USDOE Cross Section Evaluation Working Group. Upton (NY): National Nuclear Data Center, Brookhaven National Laboratory. . Retrieved 2009-01-03. [32] Fermi, Enrico (December 12, 1938). " Artificial radioactivity produced by neutron bombardment: Nobel Lecture (http:/ / www. nobel. se/ physics/ laureates/ 1938/ fermi-lecture. pdf)" (PDF). Royal Swedish Academy of Sciences. . [33] Darden, Lindley (1998). " The Nature of Scientific Inquiry (http:/ / www. philosophy. umd. edu/ Faculty/ LDarden/ sciinq/ )". College Park (MD): Department of Philosophy, University of Maryland. . Retrieved 2008-01-03. [34] LBNL contributors. " Elements 93 and 94 (http:/ / acs. lbl. gov/ Seaborg. talks/ 65th-anniv/ 14. html)". Advanced Computing for Science Department, Lawrence Berkeley National Laboratory. . Retrieved 2008-09-17. [35] This was not the first time somebody suggested that an element be named "plutonium." A decade after barium was discovered, a Cambridge University professor suggested it be renamed to "plutonium" because the element was not (as suggested by the Greek root, barys, it was named for) heavy. He reasoned that, since it was produced by the relatively new technique of electrolysis, its name should refer to fire. Thus he suggested it be named for the Roman god of the underworld, Pluto. (Heiserman 1992) [36] Clark, David L.; Hobart, David E. (2000). " Reflections on the Legacy of a Legend: Glenn T. Seaborg, 1912–1999 (http:/ / www. fas. org/ sgp/ othergov/ doe/ lanl/ pubs/ 00818011. pdf)" (PDF). Los Alamos Science 26: 56–61, on 57. . Retrieved 2009-02-15. [37] As one article puts it, referring to information Seaborg gave in a talk: "The obvious choice for the symbol would have been Pl, but facetiously, Seaborg suggested Pu, like the words a child would exclaim, 'Pee-yoo!' when smelling something bad. Seaborg thought that he would receive a great deal of flak over that suggestion, but the naming committee accepted the symbol without a word."
Clark, David L.; Hobart, David E. (2000). " Reflections on the Legacy of a Legend: Glenn T. Seaborg, 1912–1999 (http:/ / www. fas. org/ sgp/ othergov/ doe/ lanl/ pubs/ 00818011. pdf)" (PDF). Los Alamos Science 26: 56–61, on 57. . Retrieved 2009-02-15. [40] PBS contributors (1997). " Frontline interview with Seaborg (http:/ / www. pbs. org/ wgbh/ pages/ frontline/ shows/ reaction/ interviews/ seaborg. html)". Frontline. Public Broadcasting Service. . Retrieved 2008-12-07. [41] NPS contributors. " Room 405, George Herbert Jones Laboratory (http:/ / tps. cr. nps. gov/ nhl/ detail. cfm?ResourceId=735& ResourceType=Building)". National Park Service. . Retrieved 2008-12-14. [42] Room 405 of the George Herbert Jones Laboratory, where the first isolation of plutonium took place, was named a National Historic Landmark in May 1967. [43] Miner 1968, p. 540 [44] LANL contributors. " Site Selection (http:/ / www. lanl. gov/ history/ road/ siteselection. shtml)". LANL History. Los Alamos, New Mexico: Los Alamos National Laboratory. . Retrieved 2008-12-23. [45] During the Manhattan Project, plutonium was also often referred to as simply "49": the number 4 was for the last digit in 94 (atomic number of plutonium), and 9 was for the last digit in Pu-239, the weapon-grade fissile
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Plutonium isotope used in nuclear bombs.
Hammel, E.F. (2000). " The taming of "49" – Big Science in little time. Recollections of Edward F. Hammel, pp. 2-9. In: Cooper N.G. Ed. (2000). Challenges in Plutonium Science (http:/ / www. fas. org/ sgp/ othergov/ doe/ lanl/ pubs/ 00818010. pdf)". Los Alamos Science 26 (1): 2–9. . Retrieved 2009-02-15. Hecker, S.S. (2000). " Plutonium: an historical overview. In: Challenges in Plutonium Science (http:/ / www. fas. org/ sgp/ othergov/ doe/ lanl/ pubs/ number26. htm)". Los Alamos Science 26 (1): 1–2. . Retrieved 2009-02-15. [48] Sublette, Carey. " Atomic History Timeline 1942-1944 (http:/ / www. atomicheritage. org/ index. php?option=com_content& task=view& id=288& Itemid=202)". Washington (DC): Atomic Heritage Foundation. . Retrieved 2008-12-22. [49] Wahlen, R.K. (1989) (PDF). History of 100-B Area (http:/ / www. hanford. gov/ doe/ history/ files/ HistoryofBArea. pdf). Richland, Washington: Westinghouse Hanford Company. pp. iv, 1. WHC-EP-0273. . Retrieved 2009-02-15. [50] The American Society of Mechanical Engineers (ASME) established B Reactor as a National Historic Mechanical Engineering Landmark in September 1976.
Wahlen, R.K. (1989) (PDF). History of 100-B Area (http:/ / www. hanford. gov/ doe/ history/ files/ HistoryofBArea. pdf). Richland, Washington: Westinghouse Hanford Company. p. 1. WHC-EP-0273. . Retrieved 2009-02-15. [52] In August 2008, B Reactor was designated a U.S. National Historic Landmark.
" Weekly List Actions (http:/ / www. nps. gov/ history/ nr/ listings/ 20080829. HTM)". National Park Service. 2008-08-29. . Retrieved 2008-08-30. [55] Rincon, Paul (2009). " BBC NEWS – Science & Environment – US nuclear relic found in bottle (http:/ / news. bbc. co. uk/ 2/ hi/ science/ nature/ 7918618. stm)". . Retrieved 2009-03-02. [56] Gebel, Erika (2009). " Old plutonium, new tricks (http:/ / pubs. acs. org/ doi/ abs/ 10. 1021/ ac900093b)". Analytical Chemistry 81 (5): 1724. doi: 10.1021/ac900093b (http:/ / dx. doi. org/ 10. 1021/ ac900093b). . Retrieved 2009-03-02. [57] Schwantes, Jon M.; Matthew Douglas, Steven E. Bonde, James D. Briggs, Orville T. Farmer, Lawrence R. Greenwood, Elwood A. Lepel, Christopher R. Orton, John F. Wacker, Andrzej T. Luksic (2009). " Nuclear archeology in a bottle: Evidence of pre-Trinity U.S. weapons activities from a waste burial site (http:/ / dx. doi. org/ 10. 1021/ ac802286a)". Analytical Chemistry 81 (4): 1297-1306. doi: 10.1021/ac802286a (http:/ / dx. doi. org/ 10. 1021/ ac802286a). . Retrieved 2009-03-02. [58] Sublette, Carey (2007-07-03). " 8.1.1 The Design of Gadget, Fat Man, and "Joe 1" (RDS-1) (http:/ / nuclearweaponarchive. org/ Nwfaq/ Nfaq8. html#nfaq8. 1. 1)". Nuclear Weapons Frequently Asked Questions. The Nuclear Weapon Archive. . Retrieved 2008-01-04. [59] Malik, John (September 1985). The Yields of the Hiroshima and Nagasaki Explosions (http:/ / www. fas. org/ sgp/ othergov/ doe/ lanl/ docs1/ 00313791. pdf). Los Alamos. p. Table VI. LA-8819. . Retrieved 2009-02-15. [60] The efficiency calculation is based on the fact that 1 kg of Pu-239 (or U-235) fissioning results in an energy release of approximately 17 kt, leading to a rounded estimate of 1.2 kg plutonium actually fissioned to produce the 20 kt yield. On the figure of 1 kg = 17 kt,
" Proliferation of Nuclear Weapons and Materials to State and Non-State Actors: What It Means for the Future of Nuclear Power (http:/ / www. fas. org/ rlg/ PNWM_UMich. pdf)". University of Michigan Symposium. Federation of American Scientists. 2002-10-04. . Retrieved 2009-01-04. [62] DOE contributors (2001). Historic American Engineering Record: B Reactor (105-B Building) (http:/ / www. fas. org/ sgp/ othergov/ doe/ pu50yb. html#ZZ13). Richland (WA): U.S. Department of Energy. p. 110. DOE/RL-2001-16. . Retrieved 2008-12-24. [63] Cochran, Thomas B. (1997). " Safeguarding nuclear weapons-usable materials in Russia (http:/ / docs. nrdc. org/ nuclear/ nuc_06129701a_185. pdf)". International Forum on Illegal Nuclear Traffic. Washington (DC): Natural Resources Defense Council, Inc. [64] Much of this plutonium was used to make the fissionable cores of a type of thermonuclear weapon employing the Teller–Ulam design. These so-called 'hydrogen bombs' are a variety of nuclear weapon that use a fission bomb to trigger the nuclear fusion of heavy hydrogen isotopes. Their destructive yield is commonly in the millions of tons of TNT equivalent compared with the thousands of tons of TNT equivalent of fission-only
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Plutonium devices.(Emsley 2001) [65] Stockholm International Peace Research Institute (2007). SIPRI Yearbook 2007: Armaments, Disarmament, and International Security (http:/ / books. google. com/ books?id=2M0C6SERFG0C& pg=PA567). Oxford University Press. p. 567. ISBN 0199230218, 9780199230211. . [66] Gadolinium zirconium oxide (Gd2Zr2O7) has been studied because it could hold plutonium for up to 30 million years.(Emsley 2001) [67] Press Secretary (July 23, 2002). " President Signs Yucca Mountain Bill (http:/ / web. archive. org/ web/ 20080306193653/ http:/ / georgewbush-whitehouse. archives. gov/ news/ releases/ 2002/ 07/ 20020723-2. html)". Washington (DC): Office of the Press Secretary, White House. Archived from the original (http:/ / georgewbush-whitehouse. archives. gov/ news/ releases/ 2002/ 07/ 20020723-2. html) on 2008-03-06. . Retrieved 2009-01-04. [68] Moss, William; Eckhardt, Roger (1995). " The Human Plutonium Injection Experiments (http:/ / library. lanl. gov/ cgi-bin/ getfile?00326640. pdf)" (PDF). Los Alamos Science (Los Alamos National Laboratory) 23: 188, 205, 208, 214. . Retrieved 2006-06-06. [69] Voelz, George L. (2000). "Plutonium and Health: How great is the risk?". Los Alamos Science (Los Alamos (NM): Los Alamos National Laboratory) (26): 78–79. [70] Yesley, Michael S. (1995). " 'Ethical Harm' and the Plutonium Injection Experiments (http:/ / www. fas. org/ sgp/ othergov/ doe/ lanl/ pubs/ 00326649. pdf)" (PDF). Los Alamos Science 23: 280–283. . Retrieved 2009-02-15. [71] Martin, James E. (2000). Physics for Radiation Protection (1st ed.). Wiley-Interscience. p. 532. ISBN 0471353736. [72] FAS contributors (1998). " Nuclear Weapon Design (http:/ / www. fas. org/ nuke/ intro/ nuke/ design. htm)". Federation of American Scientists. . Retrieved 2008-12-07. [73] WNA contributors (2006). " Mixed Oxide Fuel (MOX) (http:/ / www. world-nuclear. org/ info/ inf29. html)". London (UK): World Nuclear Association. . Retrieved 2008-12-14. [74] DNFSB staff (2004) (PDF). Plutonium Storage at the Department of Energy's Savannah River Site: First Annual Report to Congress (http:/ / www. hss. energy. gov/ deprep/ 2004/ fb04y28b. pdf). Defense Nuclear Facilities Safety Board. pp. A–1. . Retrieved 2009-02-15. (public domain text) [75] Breakdown of plutonium in a spent nuclear fuel rod: Pu-239 (~58%), Pu-240 (24%), Pu-241 (11%), Pu-242 (5%), and Pu-238 (2%). (Emsley 2001) [76] Besmann, Theodore M. (2005). "Thermochemical Behavior of Gallium in Weapons-Material-Derived Mixed-Oxide Light Water Reactor (LWR) Fuel". Journal of the American Ceramic Society 81 (12): 3071–3076. doi: 10.1111/j.1151-2916.1998.tb02740.x (http:/ / dx. doi. org/ 10. 1111/ j. 1151-2916. 1998. tb02740. x). 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[80] Defunct pacemakers with Pu power source (http:/ / www. orau. org/ ptp/ collection/ Miscellaneous/ pacemaker. htm) [81] ORAU contributors (1974). " Plutonium Powered Pacemaker (http:/ / www. orau. org/ PTP/ collection/ Miscellaneous/ pacemaker. htm)". Oak Ridge (TN): Orau.org. . Retrieved 2008-09-12. [82] Bayles, John J.; Taylor, Douglas (1970). SEALAB III - Diver's Isotopic Swimsuit-Heater System (http:/ / oai. dtic. mil/ oai/ oai?verb=getRecord& metadataPrefix=html& identifier=AD0708680). Port Hueneme (CA): Naval Civil Engineering Lab. AD0708680. . [83] ATSDR contributors (2007). " Toxicological Profile for Plutonium, Draft for Public Comment (http:/ / www. atsdr. cdc. gov/ toxprofiles/ tp143. html)". U.S. Department of Health and Human Services, Agency for Toxic Substances and Disease Registry (ATSDR). . Retrieved 2008-05-22. [84] Cohen, Bernard L. (1985). Karl Otto Ott and Bernard I. Spinrad, eds. ed. Nuclear Energy (New York (NY): Plenum Press): 355–365. [85] WNA contributors (2008). " Plutonium (http:/ / www. world-nuclear. org/ info/ inf15. html)". London (UK): World Nuclear Association. . Retrieved 2008-05-22. [86] Brown, SC; Schonbeck MF, McClure D et al. (July 2004). " Lung cancer and internal lung doses among plutonium workers at the Rocky Flats Plant: a case-control study (http:/ / aje. oxfordjournals. org/ cgi/ content/ full/ 160/ 2/ 163)". American Journal of Epidemiology (Oxford Journals) 160 (2): 163–172. doi:
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Plutonium 10.1093/aje/kwh192 (http:/ / dx. doi. org/ 10. 1093/ aje/ kwh192). PMID 15234938. . Retrieved 2009-02-15. [87] ANL staff (2001). " ANL human health fact sheet--plutonium (http:/ / consolidationeis. doe. gov/ PDFs/ PlutoniumANLFactSheetOct2001. pdf)". Argonne National Laboratory. . Retrieved 2007-06-16. [88] DOE staff. " Radiological control technical training (http:/ / web. archive. org/ web/ 20070630190114/ http:/ / hss. energy. gov/ NuclearSafety/ techstds/ standard/ hdbk1122-04/ part9of9. pdf)". U.S. Department of Energy. Archived from the original (http:/ / hss. energy. gov/ NuclearSafety/ techstds/ standard/ hdbk1122-04/ part9of9. pdf) on 2007-06-30. . Retrieved 2008-12-14. [89] Welsome, Eileen (2000). The Plutonium Files: America's Secret Medical Experiments in the Cold War. New York (NY): Random House. p. 17. ISBN 0-385-31954-1. [90] Miner 1968, p. 546 [91] Roark, Kevin N. (2000). Criticality accidents report issued (http:/ / www. lanl. gov/ news/ index. php/ fuseaction/ home. story/ story_id/ 1054/ view/ print). Los Alamos (NM): Los Alamos National Laboratory. . Retrieved 2008-11-16. [92] LANL contributors. " Raemer Schreiber (http:/ / www. lanl. gov/ history/ people/ R_Schreiber. shtml)". Staff Biographies. Los Alamos (NM): Los Alamos National Laboratory. . Retrieved 2008-11-16. [93] McLaughlin, Thomas P.; Monahan, Shean P.; Pruvost, Norman L. (2000). A Review of Criticality Accidents. Los Alamos (NM): Los Alamos National Laboratory. p. 16. LA-13638. [94] DOE contributors. " Plutonium (http:/ / www. hss. energy. gov/ nuclearsafety/ ns/ techstds/ standard/ hdbk1081/ hbk1081d. html#ZZ281)". Nuclear Safety and the Environment. Department of Energy, Office of Health Safety and Security. . Retrieved 2008-12-07. [95] DOE contributors (1994). " Primer on Spontaneous Heating and Pyrophoricity – Pyrophoric Metals – Plutonium (http:/ / web. archive. org/ web/ 20070428220410/ http:/ / www. hss. energy. gov/ NuclearSafety/ techstds/ standard/ hdbk1081/ hbk1081d. html#ZZ281)". Washington (DC): U.S. Department of Energy, Office of Nuclear Safety, Quality Assurance and Environment. Archived from the original (http:/ / www. hss. energy. gov/ NuclearSafety/ techstds/ standard/ hdbk1081/ hbk1081d. html#ZZ281) on 2007-04-28. . [96] There was a major plutonium-initiated fire at the Rocky Flats Plant near Boulder, Colorado in 1969.
Albright, David; O'Neill, Kevin (1999). " The Lessons of Nuclear Secrecy at Rocky Flats (http:/ / www. isis-online. org/ publications/ usfacilities/ Rfpbrf. html)". ISIS Issue Brief. Institute for Science and International Security (ISIS). . Retrieved 2008-12-07.
Bibliography • CRC contributors (2006). David R. Lide. ed. Handbook of Chemistry and Physics (87th ed.). Boca Raton (FL): CRC Press, Taylor & Francis Group. ISBN 0849304873. • Emsley, John (2001). "Plutonium". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford (UK): Oxford University Press. pp. 324–329. ISBN 0198503407. • Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.). Oxford (UK): Butterworth-Heinemann. ISBN 0-7506-3365-4. • Heiserman, David L. (1992). "Element 94: Plutonium". Exploring Chemical Elements and their Compounds. New York (NY): TAB Books. pp. 337–340. ISBN 0-8306-3018-X. • Miner, William N.; Schonfeld, Fred W. (1968). "Plutonium". in Clifford A. Hampel (editor). The Encyclopedia of the Chemical Elements. New York (NY): Reinhold Book Corporation. pp. 540–546. LCCN 68-29938. • Stwertka, Albert (1998). "Plutonium". Guide to the Elements (Revised ed.). Oxford (UK): Oxford University Press. ISBN 0-19-508083-1.
External links • Sutcliffe, W.G.; et al. (1995). " A Perspective on the Dangers of Plutonium (http:/ / web. archive. org/ web/ 20060929015050/ http:/ / www. llnl. gov/ csts/ publications/ sutcliffe/ )". Lawrence Livermore National Laboratory. Archived from the original (http:/ / www. llnl. gov/ csts/ publications/ sutcliffe/ ) on 2006-09-29. http:/ / web. archive. org/ web/ 20060929015050/ http:/ / www. llnl. gov/ csts/ publications/ sutcliffe/ .
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Plutonium • Johnson, C.M.; Davis, Z.S. (1997). " Nuclear Weapons: Disposal Options for Surplus Weapons-Usable Plutonium (http:/ / www. globalsecurity. org/ wmd/ library/ report/ crs/ 97-564. htm)". CRS Report for Congress # 97-564 ENR. http:/ / www. globalsecurity. org/ wmd/ library/ report/ crs/ 97-564. htm. Retrieved 2009-02-15. • IEER contributors (2005). " Physical, Nuclear, and Chemical, Properties of Plutonium (http:/ / www. ieer. org/ fctsheet/ pu-props. html)". IEER. http:/ / www. ieer. org/ fctsheet/ pu-props. html. Retrieved 2009-02-15. • Samuels, D. (2005). " End of the Plutonium Age (http:/ / discovermagazine. com/ 2005/ nov/ end-of-plutonium)". Discover Magazine 26 (11). http:/ / discovermagazine. com/ 2005/ nov/ end-of-plutonium. • Pike, J. (2000). " Plutonium production (http:/ / www. fas. org/ nuke/ intro/ nuke/ plutonium. htm)". Federation of American Scientists. http:/ / www. fas. org/ nuke/ intro/ nuke/ plutonium. htm. Retrieved 2009-02-15. • Nuclear Weapon Archive contributors. " Plutonium Manufacture and Fabrication (http:/ / nuclearweaponarchive. org/ Library/ Plutonium/ )". Nuclearweaponarchive.org. http:/ / nuclearweaponarchive. org/ Library/ Plutonium/ . • Ong, C. (1999). " World Plutonium Inventories (http:/ / www. nuclearfiles. org/ menu/ key-issues/ nuclear-energy/ issues/ world-plutonium-inventories-ong. htm)". Nuclear Files.org. http:/ / www. nuclearfiles. org/ menu/ key-issues/ nuclear-energy/ issues/ world-plutonium-inventories-ong. htm. Retrieved 2009-02-15. • LANL contributors (2000). " Challenges in Plutonium Science (http:/ / www. fas. org/ sgp/ othergov/ doe/ lanl/ pubs/ number26. htm)". Los Alamos Science I & II (26). http:/ / www. fas. org/ sgp/ othergov/ doe/ lanl/ pubs/ number26. htm. Retrieved 2009-02-15. • NLM contributors. " Plutonium, Radioactive (http:/ / toxnet. nlm. nih. gov/ cgi-bin/ sis/ search/ r?dbs+ hsdb:@term+ @na+ @rel+ plutonium,+ radioactive)". NLM Hazardous Substances Databank. http:/ / toxnet. nlm. nih. gov/ cgi-bin/ sis/ search/ r?dbs+ hsdb:@term+ @na+ @rel+ plutonium,+ radioactive. Retrieved 2009-02-15. • Alsos contributors. " Annotated Bibliography on plutonium (http:/ / alsos. wlu. edu/ qsearch. aspx?browse=science/ Plutonium)". Alsos Digital Library for Nuclear Issues. http:/ / alsos. wlu. edu/ qsearch. aspx?browse=science/ Plutonium. Retrieved 2009-02-15. pnb:مینوٹولپ
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Americium
1
Americium plutonium ← americium → curiumEu ↑ Am ↓ (Uqp)
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
95Am Periodic table
Appearance silvery white General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Melting pointBoiling pointHeat of fusionSpecific heat capacityVapor pressure P/Pa
1
10
at T/K
1239
1356
100
1k
10 k
100 k
Atomic properties Oxidation states ElectronegativityIonization energiesAtomic radiusCovalent radius Miscellaneous Crystal structureMagnetic orderingThermal conductivityCAS registry number Most stable isotopes Main article: Isotopes of americium iso 241
Am
242m
Am
N.A. syn
syn
half-life 432.2 y
141 y
DM
DE (MeV)
DP
SF
-
-
α
5.486
237
IT
0.049
242
α
5.637
238
SF
-
-
Np Am Np
Americium
243
Am
2 syn
7370 y
SF
-
-
α
5.275
239
Np
americium, Am, 95 actiniden/a, 7, f(243) g·mol−1 [Rn] 5f7 7s2 2, 8, 18, 32, 25, 8, 2 (Image) solid 12 g·cm−3 1449 K,1176 °C,2149 °F 2880 K,2607 °C,4725 °F 14.39 kJ·mol−1 (25 °C) 62.7 J·mol−1·K−1 6, 5, 4, 3, 2 (amphoteric oxide) 1.3 (Pauling scale) 1st: 578 kJ·mol−1173 pm 180±6 pm hexagonal no data (300 K) 10 W·m−1·K−1 7440-35-9 Americium (pronounced /ˌæməˈrɪsiəm/) is a synthetic element that has the symbol Am and atomic number 95. A radioactive metallic element, americium is an actinide that was obtained in 1944 by Glenn T. Seaborg who was bombarding plutonium with neutrons and was the fourth transuranic element to be discovered. It was named for the Americas, by analogy with europium.[1] Americium is widely used in commercial ionization chamber smoke detectors, as well as in neutron sources and industrial gauges.
Properties Physical Pure americium has a silvery and white luster. At room temperatures it slowly tarnishes in dry air. It is more silvery than plutonium or neptunium and apparently more malleable than neptunium or uranium. Alpha emission from 241Am is approximately three times that of radium. Gram quantities of 241Am emit intense gamma rays which creates a serious exposure problem for anyone handling the element. Americium is also fissile; the critical mass for an unreflected sphere of 241Am is approximately 60 kilograms. It is unlikely that Americium would be used as a weapons material, as its minimum critical mass is considerably larger than that of more readily obtained plutonium or uranium isotopes.[2]
Americium
3
Chemical Americium oxidizes to AmO in air. Similarly, reaction with hydrogen results in AmH2 where Am is divalent. However, the most common oxidation state of Am is +3, especially in solutions which are colored red. It is much harder to oxidize Am(III) to Am(IV) than it is to oxidize Pu(III) to Pu(IV). Americium, unlike uranium, does not readily form a dioxide americyl core (AmO2).[3] This is because americium is very hard to oxidise above the +3 oxidation state when it is in an aqueous solution. In the environment, this americyl core could complex with carbonate as well as other oxygen moieties (OH−, NO−2, NO−3, and SO2−4) to form charged complexes which tend to be readily mobile with low affinities to soil: AmO2(OH)+, AmO2(OH)2+2, AmO2CO+3, AmO2(CO3)−2 and AmO2(CO3)3−3. Examples of americium +4 compounds are Am(OH)4 and AmF4. All pentavalent and hexavalent americium compounds are complex salts such as KAmO2F2, Li3AmO4 and Li6AmO6, Ba3AmO6, AmO2F2. Hexavalent americium is a strong oxidizing agent and is reduced to AmO2+ in oxidation-reduction reactions.[4] World's first sample of americium (as the hydroxide)
Extraction A large amount of work has been done on the solvent extraction of americium, as it is the case that americium and the other transplutonium elements are responsible for the majority of the long lived radiotoxicity of spent nuclear fuel. It is thought that by removal of the americium and curium that the used fuel will only need to be isolated from people and the environment for a shorter time than that required for the isolation of untreated used fuel. One recent EU funded project on this topic was known by the codename "EUROPART". Within this project triazines and other compounds were studied as potential extraction agents.[5] [6] [7] [8] [9]
Isotopes Eighteen radioisotopes of americium have been characterized, with the most stable being 243 Am with a half-life of 7370 years, and 241Am with a half-life of 432.2 years. All of the remaining radioactive isotopes have half-lives that are less than 51 hours, and the majority of these have half-lives that are less than 100 minutes. This element also has 8 meta states, with the most stable being 242mAm (t½ 141 years). The isotopes of americium range in atomic weight from 231.046 u (231Am) to 249.078 u (249Am).
Americium
4
History Americium was first isolated by Glenn T. Seaborg, Leon O. Morgan, Ralph A. James, and Albert Ghiorso in late 1944 at the wartime Metallurgical Laboratory at the University of Chicago (now known as Argonne National Laboratory). The team created the isotope 241Am by subjecting 239Pu to successive neutron capture reactions in a nuclear reactor. This created 240Pu and then 241Pu which in turn decayed into 241Am via beta decay.[10]
Seaborg was granted a patent for "Element 95 and Method of Producing Said Element," whose unusually terse claim number 1 reads simply, "Element 95."[11] The discovery of americium and curium was first announced informally on a children's quiz show in 1945.[12]
Applications Americium can be produced in kilogram amounts and has some uses, mostly involving 241Am since it is easiest to produce relatively pure samples of this isotope. Americium is the only synthetic element to have found its way into the household, where one common type of smoke detector uses 241Am in the form of americium dioxide as its source of ionizing radiation.[13] The amount of americium in a typical smoke detector when new is 1 microcurie or 0.28 microgram. This amount declines slowly as the americium decays into neptunium-237, a different transuranic element with a much longer half-life (about 2.14 million years). With its half-life of 432.2 years, the americium in a smoke detector includes about 3% neptunium after 19 years, and about 5% after 32 years.
Americium-241 is used in smoke detectors
241
Am has been used as a portable source of both gamma rays and alpha particles for a number of medical and industrial uses. Gamma ray emissions from 241Am can be used for indirect analysis of materials radiography and for quality control in manufacturing fixed gauges. For example, the element has been employed to gauge glass thickness to help create flat glass. 241Am gamma rays were also used to provide passive diagnosis of thyroid function. This medical application is obsolete. 241Am can be combined with lighter elements (e.g., beryllium or lithium) to become a neutron emitter. This application has found uses in neutron radiography as well as a neutron emitting radioactive source. The most widespread use of 241AmBe neutron sources is found in moisture/density gauges used for quality control in highway construction. 241Am neutron sources are also critical for well logging applications. 242mAm has been cited for use as an advanced nuclear rocket propulsion fuel.[14] [15] This isotope is, however, extremely expensive to produce in usable quantities. 241
Am has recently been suggested for use as a denaturing agent in plutonium reactor fuel rods to render the fuel unusable for conversion to nuclear weapons.[16]
Americium
5
Safety Americium emits alpha and gamma radiation. The alpha decay of 241Am is three times as active as that of radium. It is associated with 5.48 MeV alpha particles and 59 keV gamma emission, which is a serious health hazard.[4]
See also • Actinides in the environment
Further reading • Nuclides and Isotopes - 14th Edition, GE Nuclear Energy, 1989. • Gabriele Fioni, Michel Cribier and Frédéric Marie. "Can the minor actinide, [17] americium-241, be transmuted by thermal neutrons? ". Commissariat à l'énergie atomique. http:/ / www. cea. fr/ var/ cea/ storage/ static/ gb/ library/ Clefs46/ pagesg/ clefs46_30. html. • Guide to the Elements - Revised Edition, Albert Stwertka, (Oxford University Press; 1998) ISBN 0-19-508083-1 • Gmelins Handbuch der anorganischen Chemie, System Nr. 71, Band 7 a, Transurane: Teil A 1 I, S. 30–34; Teil A 1 II, S. 18, 315–326, 343–344; Teil A 2, S. 42–44, 164–175, 185–188; Teil B 1, S. 57–67.
External links • • • •
WebElements.com – Americium [18] It's Elemental – Americium [19] ATSDR – Public Health Statement: Americium - World Nuclear Association [21]
[20]
References [1] Seaborg, Glenn T. (1946). " The Transuranium Elements (http:/ / www. jstor. org/ stable/ 1675046)". Science 104 (2704): 379–386. doi: 10.1126/science.104.2704.379 (http:/ / dx. doi. org/ 10. 1126/ science. 104. 2704. 379). . [2] " Fissile Materials & Nuclear Weapons: Introduction (http:/ / www. fissilematerials. org/ ipfm/ pages_us_en/ fissile/ fissile/ fissile. php)". International Panel on Fissile Materials. . Retrieved 2007-11-22. [3] David L. Clark (2000). " The Chemical Complexities of Plutonium (http:/ / fas. org/ sgp/ othergov/ doe/ lanl/ pubs/ 00818038. pdf)" (Reprinted at fas.org). Los Alamos Science (26). . [4] Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds (http:/ / books. google. com/ books?id=Xqj-TTzkvTEC& pg=PA18). McGraw-Hill. p. 18. ISBN 0070494398. . Retrieved 2009-06-06. [5] Michael J. Hudson, Michael G. B. Drew, Mark R. StJ. Foreman, Clément Hill, Nathalie Huet, Charles Madic and Tristan G. A. Youngs (2003). "The coordination chemistry of 1,2,4-triazinyl bipyridines with lanthanide(III) elements – implications for the partitioning of americium(III)". Dalton Trans.: 1675–1685. doi: 10.1039/b301178j (http:/ / dx. doi. org/ 10. 1039/ b301178j). [6] Andreas Geist, Michael Weigl, Udo Müllich, Klaus Gompper (11-13 December 2000). " Actinide(III)/Lanthanide(III) Partitioning Using n-Pr-BTP as Extractant: Extraction Kinetics and Extraction Test in a Hollow Fiber Module (http:/ / www. nea. fr/ html/ pt/ docs/ iem/ madrid00/ Paper14. pdf)" (PDF). 6th Information Exchange Meeting on Actinide and Fission Product Partitioning and Transmutation. OECD Nuclear Energy Agency. . [7] C. Hill, D. Guillaneux, X. Hérès, N. Boubals and L. Ramain (24-26 October 2000). " Sanex-BTP Process Development Studies (http:/ / www-atalante2004. cea. fr/ home/ liblocal/ docs/ atalante2000/ P3-26. pdf)" (PDF). Atalante 2000: Scientific Research on the Back-end of the Fuel Cycle for the 21st Century. Commissariat à l'énergie atomique. .
Americium [8] Andreas Geist, Michael Weigl and Klaus Gompper (14-16 October 2002). " Effective Actinide(III)-Lanthanide(III) Separation in Miniature Hollow Fibre Modules (http:/ / www. nea. fr/ html/ pt/ docs/ iem/ jeju02/ session2/ SessionII-15. pdf)" (PDF). 7th Information Exchange Meeting on Actinide and Fission Product Partitioning and Transmutation. OECD Nuclear Energy Agency. . [9] D.D. Ensor. " Separation Studies of f-Elements (http:/ / www. tntech. edu/ WRC/ pdfs/ Projects04_05/ Ens_Elem. pdf)" (PDF). Tennessee Tech University. . [10] G. T. Seaborg, R. A. James, L. O. Morgan: "The New Element Americium (Atomic Number 95)", NNES PPR (National Nuclear Energy Series, Plutonium Project Record), Vol. 14 B The Transuranium Elements: Research Papers, Paper No. 22.1, McGraw-Hill Book Co., Inc., New York, 1949; Abstract (http:/ / www. osti. gov/ cgi-bin/ rd_accomplishments/ display_biblio. cgi?id=ACC0046& numPages=43& fp=N); Typoskript (Januar 1948) (http:/ / www. osti. gov/ accomplishments/ documents/ fullText/ ACC0046. pdf). [11] Patent US3156523 (http:/ / patft. uspto. gov/ netacgi/ nph-Parser?patentnumber=3156523) (1964-11-10) Glenn T. Seaborg, Element 95 and Method of Producing Said Element. [12] Rachel Sheremeta Pepling. " It's Elemental: The Periodic Table: Americium (http:/ / pubs. acs. org/ cen/ 80th/ americium. html)". Chemical & Engineering News. . [13] Americium dioxide is used in smoke detectors. (Internet Archive) (http:/ / web. archive. org/ web/ 19960101-re_/ http:/ / www. uic. com. au/ nip35. htm) [14] " Extremely Efficient Nuclear Fuel Could Take Man To Mars In Just Two Weeks (http:/ / www. sciencedaily. com/ releases/ 2001/ 01/ 010103073253. htm)". ScienceDaily. 2001-01-03. . Retrieved 2007-11-22. [15] Terry Kammash, David L. Galbraith, and Ta-Rong Jan (January 10, 1993). "An americium-fueled gas core nuclear rocket". Tenth symposium on space nuclear power and propulsion. AIP Conf. Proc.. 271. pp. 585-589. doi: 10.1063/1.43073 (http:/ / dx. doi. org/ 10. 1063/ 1. 43073). [16] " BGU combats nuclear proliferation (http:/ / www. jpost. com/ servlet/ Satellite?cid=1235898328437& pagename=JPost/ JPArticle/ ShowFull)". . Retrieved 2009-03-05. [17] [18] [19] [20] [21]
http:/ / www. cea. fr/ var/ cea/ storage/ static/ gb/ library/ Clefs46/ pagesg/ clefs46_30. html http:/ / www. webelements. com/ webelements/ elements/ text/ Am/ index. html http:/ / education. jlab. org/ itselemental/ ele095. html http:/ / www. atsdr. cdc. gov/ toxprofiles/ phs156. html http:/ / world-nuclear. org/ info/ inf57. html
6
Article Sources and Contributors
Article Sources and Contributors Americium Source: http://en.wikipedia.org/w/index.php?oldid=309056582 Contributors: AWeishaupt, Abeg92, Ahoerstemeier, AlimanRuna, Andres, Angela, Ayengar, B, Baccyak4H, Bayou Banjo, Bcorr, Beetstra, Bobo The Ninja, Bogey97, Bovineone, Bryan Derksen, Burtonpe, Cadmium, Camw, Canthusus, Capricorn42, CaptainVindaloo, ChemGardener, ChemNerd, Chemkid1, Conversion script, Crazytales, DV8 2XL, Dachshund, Darrien, David Latapie, David.Mestel, Davidprior, Deor, Dispenser, DocWatson42, Donarreiskoffer, Dspradau, Edgar181, Emperorbma, Espi, Farseer, Femto, Flehmen, ForestAngel, Glc9144, Glenn, Gluck 123, Goldenband, Gtstricky, Hak-kâ-ngìn, HazyM, Hemmingsen, Herbee, Hqb, ICAPTCHA, Icairns, Ideyal, Ishikawa Minoru, Itub, J miester25, J.delanoy, JWB, JWBE, Jaraalbe, Jiang, Jimfbleak, Joanjoc, JoaoRicardo, John, Julesd, Kalamkaar, Kay Dekker, Keilana, Kelovy, Kurykh, Kwamikagami, LarryMorseDCOhio, Likeitsmyjob, LindsayH, Lord Voldemort, Lottiotta, Luna Santin, Magnus Manske, Marc Venot, Materialscientist, Mav, Mdf, Murtasa, Nergaal, Nick Y., No1lakersfan, Nofutureuk, Panu, Perlmonger42, Pharaoh of the Wizards, Polonium, Poolkris, Potatoswatter, Pras, Pseudomonas, Rcnet, Recognizance, Redux, Remember, Reza kalani, Rifleman 82, Rjwilmsi, Roberta F., Robyrockets, Roentgenium111, Romanm, Samw, Saperaud, Sargentzan, Schneelocke, Semperf, Sengkang, Sheila Rogers, Sionus, Skepicalcynic, Sl, Socrates2008, Stifynsemons, Stone, Stratocracy, StuartH, Supremeknowledge, TJRC, Tagishsimon, Tetracube, The Firewall, Tim Q. Wells, Tlusťa, Tmopkisn, Tsogo3, Ttiotsw, Ttony21, Uwe W., Vary, Vsmith, WODUP, WRK, Warut, WeniWidiWiki, Wmahan, Wrynne, Xenophon777, Xxis, Yyy, Zelmerszoetrop, 190 anonymous edits
Image Sources, Licenses and Contributors file:hexagonal.svg Source: http://en.wikipedia.org/w/index.php?title=File:Hexagonal.svg License: BSD Contributors: Original uploader was Danieljamesscott at en.wikipedia file:Electron shell 095 Americium.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_095_Americium.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 Image:Americium.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Americium.jpg License: Public Domain Contributors: Berkeley-Laboratory Image:Residential smoke detector.jpg Source: http://en.wikipedia.org/w/index.php?title=File:Residential_smoke_detector.jpg License: Public Domain Contributors: User:Oleg Alexandrov
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
7
Curium
1
Curium This article is about the chemical element Curium; for the ancient city also called Curium (located in Cyprus), see Kourion
americium ← curium → berkelium Gd ↑ Cm ↓ (Uqh) Periodic Table Extended Periodic Table General Name, symbol, number
curium, Cm, 96
Element category
actinides
Group, Period, Block
n/a, 7, f
Appearance
silvery Standard atomic weight
−1
(247) g·mol
Electron configuration
7
1
2
[Rn] 5f 6d 7s
Electrons per shell
2, 8, 18, 32, 25, 9, 2 Physical properties
Phase
solid Density (near r.t.)
−3
13.51 g·cm
Melting point
1613 K (1340 °C, 2444 °F) Boiling point
3383 K (3110 °C, 5630 °F)
Heat of fusion
? 15 kJ·mol−1 Vapor pressure
P(Pa)
1
10
at T(K)
1788
1982
100
1k
10 k
Atomic properties Crystal structure Oxidation states Electronegativity Ionization energies
hexagonal close-packed 4, 3 (amphoteric oxide) 1.3 (Pauling scale) 1st: 581 kJ/mol
100 k
Curium
2 Atomic radius
174 pm
Covalent radius
169±3 pm Miscellaneous
Magnetic ordering
no data
CAS registry number
7440-51-9 Most-stable isotopes
Main article: Isotopes of curium iso 242
Cm
243
Cm
244
Cm
245
Cm
246
Cm
NA syn
syn
syn
syn
syn
half-life 160 days
α
6.1
238
α
6.169
239
ε
0.009
243
SF
-
-
SF
-
-
α
5.902
SF
-
α
5.623
241
α
5.475
242
SF
-
7
α
5.353
243
5
α
5.162
244
SF
-
-
SF
-
-
α
5.169
246
β−
0.037
250
29.1 y
18.1 y
8500 y
4730 y
1.56×10 y
248
syn
3.40×10 y
250
Cm
syn
9000 y
DP
-
syn
Cm
DE (MeV)
SF
247
Cm
DM
Pu Pu Am
240
Pu
Pu Pu
Pu Pu
Pu Bk
References
Curium (pronounced /ˈkjuːriəm/) is a synthetic chemical element with the symbol Cm and atomic number 96. A radioactive metallic transuranic element of the actinide series, curium is produced by bombarding plutonium with alpha particles (helium ions) and was named for Marie Curie and her husband Pierre.
Curium
3
Characteristics The isotope curium-248 has been synthesized only in milligram quantities, but curium-242 and curium-244 are made in multigram amounts, which allows for the determination of some of the element's properties. Curium-244 can be made in quantity by subjecting plutonium to neutron bombardment. Curium does not occur in nature. There are few commercial applications for curium but it may one day be useful in radioisotope thermoelectric generators. Curium bio-accumulates in bone tissue where its radiation destroys bone marrow and thus stops red blood cell creation. A rare earth homolog, curium is somewhat chemically similar to gadolinium but with a more complex crystal structure. Chemically reactive, its metal is silvery-white in color and the element is more electropositive than aluminium (most trivalent curium compounds are slightly yellow). Curium has been studied greatly as a potential fuel for radioisotope thermoelectric generators (RTG). Curium-242 can generate up to 120 watts of thermal energy per gram (W/g); however, its very short half-life makes it undesirable as a power source for long-term use. Curium-242 can decay by alpha emission to plutonium-238 which is the most common fuel for RTGs. Curium-244 has also been studied as an energy source for RTGs having a maximum energy density ~3 W/g,[1] but produces a large amount of neutron radiation from spontaneous fission. Curium-243 with a ~30 year half-life and good energy density of ~1.6 W/g would seem to make an ideal fuel, but it produces significant amounts of gamma and beta radiation from radioactive decay products.
Compounds Some compounds are: • • • • • •
curium curium curium curium curium curium
dioxide (CmO2) trioxide (Cm2O3) bromide (CmBr3) chloride (CmCl3) tetrafluoride (CmF4) iodide (CmI3)
History Curium was first synthesized at the University of California, Berkeley by Glenn T. Seaborg, Ralph A. James, and Albert Ghiorso in 1944.[2] The team named the new element after Marie Curie and her husband Pierre who are famous for discovering radium and for their work in radioactivity. It was chemically identified at the Metallurgical Laboratory (now Argonne National Laboratory) at the University of Chicago. It was actually the third transuranium element to be discovered even though it is the fourth in the series. Curium-242 (half-life 163 days) and one free neutron were made by bombarding alpha particles onto a plutonium-239 target in the 60-inch cyclotron at Berkeley.[3] 23994Pu
+
42He
→
24296Cm
+
10n
Due to the fact that the discovery of the new elements, curium and americium, was closely related to the Manhattan Project the results were confidential and publication was impossible. Seaborg announced the discovery of the elements on the radio show for kids, the Quiz Kids, five days before the official presentation at an American Chemical Society
Curium
4
meeting on November 11, 1945.[4] Seaborg also patented the synthesis of the new elements.[5] Louis Werner and Isadore Perlman created a visible sample of curium-242 hydroxide at the University of California in 1947 by bombarding americium-241 with neutrons.[6] Curium was made in its elemental form in 1951 for the first time.[7] [8]
Isotopes 19 radioisotopes of curium have been characterized, with the most stable being Cm-247 with a half-life of 1.56 × 107 years, Cm-248 with a half-life of 3.40 × 105 years, Cm-250 with a half-life of 9000 years, and Cm-245 with a half-life of 8500 years. All of the remaining radioactive isotopes have half-lives that are less than 30 years, and the majority of these have half-lives that are less than 33 days. This element also has 4 meta states, with the most stable being Cm-244m (t½ 34 ms). The isotopes of curium range in atomic weight from 233.051 u (Cm-233) to 252.085 u (Cm-252).
Nuclear fuel cycle
[9] Transmutation flow between 238Pu and 244Cm in LWR. Fission percentage is 100 minus shown percentages. Total rate of transmutation varies greatly by nuclide. 245 Cm–248Cm are long-lived with negligible decay. Thermal neutron cross sections 242
243
244
245
246
247
Fission
5
617
1.04
2145
0.14
81.90
Capture
16
130
15.20
369
1.22
57
C/F ratio
3.20
0.21
14.62
0.17
8.71
0.70
Cm
Cm
Cm
Cm
Cm
Cm
Curium
5 [10]
LEU spent fuel 20 years after 53 MWd/kg burnup 3 common isotopes
51
3700
390 [11]
Fast reactor MOX fuel (avg 5 samples, burnup 66-120GWd/t) Total curium 3.09 × 10-3%
27.64%
70.16%
2.166%
0.0376%
0.000928%
The odd-mass number isotopes are fissile, the even-mass number isotopes are not and can only capture neutrons, but very slowly. Therefore in a thermal reactor the even-mass isotopes accumulate as burnup increases. The MOX which is to be used in power reactors should contain little or no curium as the neutron activation of 248Cm will create californium which is a strong neutron emitter. The californium would pollute the back end of the fuel cycle and increase the dose to workers. Hence if the minor actinides are to be used as fuel in a thermal neutron reactor, the curium should be excluded from the fuel or placed in special fuel rods where it is the only actinide present.
Applications The Curium isotopes 244Cm and 242Cm are strong alpha emitters with a halflife in the months to years range and produce considerable heat during this process. These properties make them useful for applications as alpha particle source and as heat generator in radioisotope thermoelectric generators (RTG). A 244Curium source is used for the Alpha particle X-ray spectrometer on board several American and European space missions, for example the Mars Exploration Rover[12] and the Rosetta/Philae. The use in RTG is proposed for several future missions.[13] [14]
Literature • Guide to the Elements - Revised Edition, Albert Stwertka, (Oxford University Press; 1998) ISBN 0-19-508083-1.
External links • • • • •
Los Alamos National Laboratory - Curium [15] It's Elemental – Curium [16] Human Health Fact Sheet on Curium [17] WebElements.com – Curium [18] NLM Hazardous Substances Databank – Curium, Radioactive
[19]
References [1] Gmelins Handbuch der anorganischen Chemie, System Nr. 71, Band 7 a, Transurane, Teil A 2, p. 289. [2] Hall, Nina (2000). The New Chemistry: A Showcase for Modern Chemistry and Its Applications (http:/ / books. google. com/ books?id=U4rnzH9QbT4C). Cambridge University Press. pp. 8–9. ISBN 9780521452243. . [3] G. T. Seaborg, R. A. James, A. Ghiorso: "The New Element Curium (Atomic Number 96)", NNES PPR (National Nuclear Energy Series, Plutonium Project Record), Vol. 14 B, The Transuranium Elements: Research Papers, Paper No. 22.2, McGraw-Hill Book Co., Inc., New York, 1949; Abstract (http:/ / www. osti. gov/ cgi-bin/ rd_accomplishments/ display_biblio. cgi?id=ACC0049& numPages=13& fp=N); Typoskript (January 1948) (http:/ / www. osti. gov/ accomplishments/ documents/ fullText/ ACC0049. pdf).
Curium
6
[4] PEPLING, RACHEL SHEREMETA (2003). " Chemical & Engineering News: It's Elemental: The Periodic Table – Americium (http:/ / pubs. acs. org/ cen/ 80th/ americium. html)". . Retrieved 07-12-2008. [5] Patent Nr. 3161462 bei Google Patents (http:/ / www. google. de/ patents?id=MSRXAAAAEBAJ& dq=3161462). [6] L. B. Werner, I. Perlman: "Isolation of Curium", NNES PPR (National Nuclear Energy Series, Plutonium Project Record), Vol. 14 B, The Transuranium Elements: Research Papers, Paper No. 22.5, McGraw-Hill Book Co., Inc., New York, 1949. [7] Wallmann, J. C.; Crane, W. W. T.; Cunningham, B. B. (1951). "The Preparation and Some Properties of Curium Metal". Journal of the American Chemical Society 73 (1): 493–494. doi: 10.1021/ja01145a537 (http:/ / dx. doi. org/ 10. 1021/ ja01145a537). [8] Werner, L. B., L. B. last =Werner; Perlman, I. (1951). "First Isolation of Curium"". Journal of the American Chemical Society 73 (1): 5215–5217. doi: 10.1021/ja01155a063 (http:/ / dx. doi. org/ 10. 1021/ ja01155a063). [9] Sasahara, Akihiro (April 2004). " Neutron and Gamma Ray Source Evaluation of LWR High Burn-up UO2 and MOX Spent Fuels (http:/ / www. jstage. jst. go. jp/ article/ jnst/ 41/ 4/ 448/ _pdf)". Journal of NUCLEAR SCIENCE and TECHNOLOGY 41 (4): 448–456. doi: 10.3327/jnst.41.448 (http:/ / dx. doi. org/ 10. 3327/ jnst. 41. 448). . [10] " Limited Proliferation-Resistance Benefits from Recycling Unseparated Transuranics and Lanthanides from Light-Water Reactor Spent Fuel (http:/ / www. princeton. edu/ ~globsec/ publications/ pdf/ 13_3 Kang vonhippel. pdf)" (PDF). p. 4. . [11] " Analysis of Curium Isotopes in Mixed Oxide Fuel Irradiated in Fast Reactor (http:/ / wwwsoc. nii. ac. jp/ aesj/ publication/ JNST2001/ No. 10/ 38_912-914. pdf)" (PDF). . [12] R. Rieder, R. Gellert, J. Brückner, G. Klingelhöfer, G. Dreibus, A. Yen, S. W. Squyres (2003). "The new Athena alpha particle X-ray spectrometer for the Mars Exploration Rovers". J. Geophysical Research 108: 8066. doi: 10.1029/2003JE002150 (http:/ / dx. doi. org/ 10. 1029/ 2003JE002150). [13] Miskolczy, G.; Lieb, D.P. (1990). "Radioisotope Thermionic Converters for Space Applications". Energy Conversion Engineering Conference (IECEC-90) 1: 222–226. doi: 10.1109/IECEC.1990.716874 (http:/ / dx. doi. org/ 10. 1109/ IECEC. 1990. 716874). [14] O’Brien,, R.C.; Ambrosi, R. M.; Bannister, N.P.; Howe S.D.; Atkinso, H. V. (2008). "Safe radioisotope thermoelectric generators and heat sources for space applications". Journal of Nuclear Materials 377 (3): 506–521. doi: 10.1016/j.jnucmat.2008.04.009 (http:/ / dx. doi. org/ 10. 1016/ j. jnucmat. 2008. 04. 009). [15] [16] [17] [18] [19]
http:/ / periodic. lanl. gov/ elements/ 96. html http:/ / education. jlab. org/ itselemental/ ele096. html http:/ / www. ead. anl. gov/ pub/ doc/ curium. pdf http:/ / www. webelements. com/ webelements/ elements/ text/ Cm/ index. html http:/ / toxnet. nlm. nih. gov/ cgi-bin/ sis/ search/ r?dbs+ hsdb:@term+ @na+ @rel+ curium,+ radioactive
Article Sources and Contributors
Article Sources and Contributors Curium Source: http://en.wikipedia.org/w/index.php?oldid=308528858 Contributors: Ahoerstemeier, AlimanRuna, Andres, Arkuat, Auric, B.d.mills, Beetstra, Benbest, BerserkerBen, BlueEarth, Borislav Dopudja, Bryan Derksen, Cadmium, Carlossuarez46, Carnildo, CommonsDelinker, Conversion script, Daniel Case, Danielle dk, Darrien, David Latapie, Deor, El C, Emerson7, Emperorbma, Enok Walker, Epbr123, Falcanary, Femto, Fibonacci, Flewis, Gaius Cornelius, Glenn4pr, Go-in, GregorB, Gökhan, Hak-kâ-ngìn, Hashar, HazyM, Hqb, IForgotToEatBreakFast, Icairns, Ideyal, J.delanoy, JTN, JWB, JWBE, Jaraalbe, Jennavecia, Jiang, Joanjoc, Karelj, Kelovy, Kipala, Knowledgeum, Kpalion, Kurykh, Kwamikagami, LA2, Marc Venot, Mav, Mentifisto, Miss Madeline, Nergaal, New4325, Nihiltres, Nosebud, Omicronpersei8, Oo64eva, Polonium, Poolkris, Pras, Pretzelpaws, Res2216firestar, Reyk, Rjwilmsi, Roberta F., Romanm, Sakus, Saperaud, Schneelocke, Sengkang, Sfuerst, Sl, Squids and Chips, Stephenb, Steve Hart, Stifynsemons, Stone, Stratocracy, Suisui, Tagishsimon, Tetracube, Thinghy, Uannis, UkPaolo, Unara, Volland, Vsmith, Warut, Watch37264, Yekrats, Yyy, 55 anonymous edits
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7
Berkelium
1
Berkelium curium ← berkelium → californiumTb ↑ Bk ↓ (Uqs)
WARNING: Table could not be rendered - ouputting plain text. Potential causes of the problem are: (a) table contains a cell with content that does not fit on a single page (b) nested tables (c) table is too wide
HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
97Bk Periodic table
Appearance silvery General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Density (near r.t.) Melting point Atomic properties Oxidation states ElectronegativityIonization energiesAtomic radius Miscellaneous Crystal structureMagnetic orderingThermal conductivityCAS registry number Most stable isotopes Main article: Isotopes of berkelium iso 245
Bk
246
Bk
N.A. syn
syn
half-life 4.94 d
1.8 d
DM
DE (MeV)
DP
ε
0.810
245
α
6.455
241
α
6.070
242
ε
1.350
246
Cm Am Am Cm
247
syn
1380 y
α
5.889
243
248
syn
>9 y
α
5.803
244
Bk Bk
Am Am
Berkelium
249
Bk
2 syn
330 d
α
5.526
245
SF
-
-
β−
0.125
249
Am
Cf
berkelium, Bk, 97 actiniden/a, 7, f(247) g·mol−1 [Rn] 5f9 7s2 2, 8, 18, 32, 27, 8, 2 (Image) solid (alpha) 14.78 g·cm−3 (beta) 13.25 g·cm−3 (beta) 1259 K,986 °C,1807 °F3, 4 1.3 (Pauling scale) 1st: 601 kJ·mol−1 170 pm hexagonal close-packed no data (300 K) 10 W·m−1·K−1 7440-40-6 Berkelium (pronounced /bərˈkiːli.əm/, less commonly English pronunciation: /ˈbɜrkli.əm/) is a synthetic element with the symbol Bk and atomic number 97. A radioactive metallic element in the actinide series, berkelium was first synthesized by bombarding americium with alpha particles (helium ions) and was named after the University of California, Berkeley. Berkelium was the fifth transuranic element to be synthesized.
Notable characteristics Weighable amounts of
249
Bk (half-life 330 days) make it
possible to determine some of its properties using macroscopic quantities. It is a silvery metal that would easily oxidize in air at elevated temperatures and would be soluble in dilute mineral acids. X-ray diffraction techniques have been used to identify various berkelium compounds such as berkelium dioxide (BkO2), berkelium fluoride (BkF3), berkelium oxychloride (BkOCl), and berkelium trioxide (BkO3).[1] In 1962 visible amounts of berkelium chloride (BkCl3) were isolated that weighed 3 billionths of a gram. The first time visible amounts of a pure berkelium compound were produced in 1958.[2]
University of California, Berkeley
Like other actinides, berkelium bio-accumulates in skeletal tissue. This element has no known uses outside of basic research and plays no biological role.
History 60-Inch-Cyclotron
Berkelium was first synthesized by Glenn T. Seaborg, Albert Ghiorso, Stanley G. Thompson, and Kenneth Street, Jr. at the University of California, Berkeley in December 1949. The team used a cyclotron to bombard a milligram-sized target of 241Am with alpha particles to produce 243 Bk (half-life 4.5 hours) and two free neutrons.[3] [4] [5] [6] 24195Am
+
42He
→
24397Bk
+2
10n
One of the longest lived isotopes of the element, 249Bk (half-life 330 days), was later synthesized by subjecting a 244Cm target to an intense beam of neutrons.
Berkelium
3
Isotopes 19 radioisotopes of berkelium have been characterized, with the most stable being 247Bk with a half-life of 1380 years, 248Bk with a half-life of >9 years, and 249Bk with a half-life of 330 days. All of the remaining radioactive isotopes have half-lives that are less than 5 days, and the majority of these have half-lives that are less than 5 hours. This element also has 2 meta states, with the most stable being 248mBk (t½ 23.7 hours). The isotopes of berkelium range in atomic weight from 235.057 u (235Bk) to 254.091 u (254Bk).
Nuclear fuel cycle In the nuclear fuel cycle, berkelium is produced by beta decay of curium. The first curium isotope to undergo beta decay is Cm-249 with a half-life of just over an hour, so Bk-249 is the only isotope of berkelium produced in significant quantities in nuclear reactors. Production of Bk-249 requires 11 successive neutron captures on uranium-238 without nuclear fission or alpha decay, so it is only produced in small amounts. 249
Bk has a moderately large neutron capture cross section of 710 barns for thermal neutrons, 1200 barns resonance integral, but very low fission cross section for thermal neutrons. If still in a thermal reactor, much of it will therefore be converted to 250Bk which quickly decays to californium-250, but some alpha decays to curium-245.
Toxicity Berkelium accumulates in the skeletal system. The radiation can cause damage to red blood cells. The maximum permissible body burden reported for the isotope Bk–249 in the human skeleton is 0.4 ng. [7]
External links • WebElements.com - Berkelium [8] • Los Alamos National Laboratory - Berkelium • It's Elemental - Berkelium [10]
[9]
• Berkeley Science Review - An Elementary Problem
[11]
References [1] "The Solution Absorption Spectrum of Bk3+ and the Crystallography of Berkelium Dioxide, Sesquioxide, Trichloride, Oxychloride, and Trifluoride", Ph.D. Thesis, Joseph Richard Peterson, October 1967, U. S. Atomic Energy Commission Document Number UCRL-17875(1967). [2] S. G. Thompson, B. B. Cunningham: "First Macroscopic Observations of the Chemical Properties of Berkelium and Californium", supplement to Paper P/825 presented at the Second Intl. Conf., Peaceful Uses Atomic Energy, Geneva, 1958. [3] S. G. Thompson, A. Ghiorso, G. T. Seaborg: "Element 97", Physical Review 1950, 77 (6), 838–839; doi: 10.1103/PhysRev.77.838.2 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 77. 838. 2). [4] S. G. Thompson, A. Ghiorso, G. T. Seaborg: "The New Element Berkelium (Atomic Number 97)", Physical Review 1950, 80 (5), 781–789; doi: 10.1103/PhysRev.80.781 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 80. 781); Abstract (http:/ / www. osti. gov/ cgi-bin/ rd_accomplishments/ display_biblio. cgi?id=ACC0045& numPages=38& fp=N); Typoscript (26. April 1950) (http:/ / www. osti. gov/ accomplishments/ documents/ fullText/ ACC0045. pdf). [5] Stanley G. Thompson, Glenn T. Seaborg: "Chemical Properties of Berkelium"; doi: 10.2172/932812 (http:/ / dx. doi. org/ 10. 2172/ 932812); Abstract (http:/ / www. osti. gov/ energycitations/ product. biblio. jsp?query_id=0& page=0& osti_id=932812); Typoscript (24. February 1950) (http:/ / www. osti. gov/ bridge/ servlets/ purl/ 932812-Rk9Mcq/ 932812. PDF).
Berkelium [6] S. G. Thompson, B. B. Cunningham, G. T. Seaborg: "Chemical Properties of Berkelium", J. Am. Chem. Soc. 1950, 72 (6), 2798–2801; doi: 10.1021/ja01162a538 (http:/ / dx. doi. org/ 10. 1021/ ja01162a538). [7] Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0070494398 [8] http:/ / www. webelements. com/ webelements/ elements/ text/ Bk/ index. html [9] http:/ / periodic. lanl. gov/ elements/ 97. html [10] http:/ / education. jlab. org/ itselemental/ ele097. html [11] http:/ / sciencereview. berkeley. edu/ articles. php?issue=1& article=elementary
4
Article Sources and Contributors
Article Sources and Contributors Berkelium Source: http://en.wikipedia.org/w/index.php?oldid=308720068 Contributors: Ahoerstemeier, Alansohn, AlimanRuna, Andres, Angela, Bart133, Beetstra, Benbest, BlueEarth, Brian Huffman, Bryan Derksen, CMW275, Carnildo, CommonsDelinker, Conversion script, Cryptoid, Darrien, Dave6, David Latapie, Doenut1793, Doulos Christos, Edgar181, El C, Emperorbma, Falcanary, Femto, FruitMart07, Gilliam, GrahamN, Hahaha2006, Hashar, Hede2000, Hqb, Ideyal, InfoCan, Ioeth, Irishguy, IvanLanin, JWB, JWBE, Jerem43, Jiang, Joanjoc, Kalamkaar, Kaobear, Karelj, Kelovy, Kurykh, Kwamikagami, Kyoko, LA2, M0th3rT3r3asa, Marc Venot, Mav, Mdf, Mygerardromance, Nergaal, Nick Y., No1lakersfan, Oo64eva, PP Jewel, PamD, Pepspirit, Polonium, Poolkris, Remember, Rettetast, Reza kalani, Roberta F., Romanm, Saperaud, Schneelocke, Scwlong, Sl, Stifynsemons, Stratocracy, Tagishsimon, Tetracube, Thricecube, VASANTH S.N., Vsmith, Wadester16, Watch37264, Whitepaw, Wimt, Xy7, Yekrats, Yyy, 89 anonymous edits
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5
Californium
1
Californium berkelium ← californium → einsteinium Dy ↑ Cf ↓ (Uqo) Periodic Table Extended Periodic Table General Name, symbol, number Element category Group, Period, Block Appearance
californium, Cf, 98 actinides n/a, 7, f silvery
Standard atomic weight Electron configuration Electrons per shell
−1
(251) g·mol 10
[Rn] 5f
2
7s
2, 8, 18, 32, 28, 8, 2 Physical properties
Phase
solid Density (near r.t.)
Melting point
15.1 g·cm
−3
1173 K (900 °C, 1652 °F) Boiling point
1743 K (1470 °C, 2678 °F) Atomic properties
Oxidation states Electronegativity Ionization energies
2, 3, 4 1.3 (Pauling scale) 1st: 608 kJ/mol Miscellaneous
CAS registry number
7440-71-3 Most-stable isotopes
Californium
2
Main article: Isotopes of californium iso 248
Cf
249
Cf
250
Cf
NA syn
252
syn
Cf
253
Cf
254
Cf
351 y
syn
syn
Cf
DM
333.5 d
syn
251
half-life
13.08 y
898 y 2.645 y
syn
17.81 d
syn
SF
-
α
6.361
DP 244
Cm
SF
-
α
6.295
245
α
6.128
246
SF
-
α
6.176
247
α
6.217
248
SF
-
Cm Cm
Cm Cm
-
0.285
253
α
6.124
249
SF
-
α
5.926
−
β
60.5 d
DE (MeV)
Es Cm
250
Cm
References
Californium (pronounced /ˌkælɨˈfɔrniəm/) is a metallic chemical element with the symbol Cf and atomic number 98. A radioactive transuranic element, californium is used in starting nuclear reactors, optimizing coal-fired power plants and cement production facilities (via online analyzers), medical treatment of cancer, and oil exploration via down hole well logging. It was first produced by bombarding curium with alpha particles (helium ions).
History Californium was first synthesized at the University of California, Berkeley by researchers Stanley G. Thompson, Kenneth Street, Jr., Albert Ghiorso and Glenn T. Seaborg in 1950. It was the sixth transuranium element to be discovered and the team announced their discovery on March 17, 1950.[1] [2] [3] They named it after the U.S. state of California and the University of California, Berkeley. Unlike the names for the elements 95 to 97, this name did not reflect the chemical homology of element 98 to its corresponding sixth-period element, dysprosium (No. 66).[2]
60 Inch Cyclotron
To produce element 98, the team bombarded a microgram-sized target of 242Cm with 35 MeV alpha particles in the 60-inch (1.52 m) Berkeley cyclotron, which produced atoms of 245 Cf (half-life 44 minutes) and a free neutron. 24296Cm
+
42He
→
24598Cf
+
10n
Californium
3 Weighable quantities of califorinium were first produced by long-duration irradiation of plutonium targets at the Materials Testing Reactor.[4] The high spontaneous fission rate of Cf-252 was observed in these samples. In the 1960s, reactors at Savannah River and the High Flux Isotope Reactor(HFIR) started producing batches of californium regularly. The U.S. Atomic Energy Commission began selling and loaning californium sources to industrial and academic customers in 1970 for $10 per microgram. By the 1990s, Oak Ridge was producing 300-400 mg of Cf, mostly Cf-252, every two years.
Milligram-quantities of californium can only be made in specialized high-flux reactors; there are only two reactors operating that can efficiently produce it, the Elutioncurves: High Flux Isotope Reactor in the U.S. and Research chromatographic separation of Dy, Tb, Institute of Atomic Reactors in Dimitrovgrad, Russia, [2] Gd, Eu and Cf, Bk, Cm, Am. with HFIR filling about 2/3 of the world market of about 90 mg / year. Between 1960 and 1995, the HFIR produced only 8 grams of californium, peaking at about 200 mg per year.[5] Plutonium supplied by the United Kingdom to the U.S. under the 1958 US-UK Mutual Defence Agreement was used for californium production.[6]
Isotopes Twenty radioisotopes of californium have been characterized, the most stable being 251Cf with a half-life of 898 years, 249Cf with a half-life of 351 years, and 250Cf with a half-life of 13 years. All of the remaining radioactive isotopes have half-lives that are less than 2.7 years, and the majority of these have half-lives shorter than 20 minutes. The isotopes of californium range in atomic weight from 237.062 u (237Cf) to 256.093 u (256Cf).
Californium
4
252
Cf has a half life of 2.645 years. 252Cf undergoes α-decay 96.9% of the time while the remaining 3.1% of decays are spontaneous fission. Each spontaneous fission decay emits an average of 3.77 neutrons per 254 fission. Cf decays nearly quantitatively by spontaneous fission with a half-life of 60.5 days. Both materials can be used as a neutron source.
Occurrence Natural occurrence Although californium does not occur naturally on Earth, the element and its decay products occur elsewhere in the universe. Their electromagnetic emissions are regularly observed in the spectra of supernovae.[8] [9]
Energy spectrum of neutrons emitted [7] by 252Cf.
[10] [11]
Fallout On November 1, 1952, the fallout of the United States hydrogen bomb test Ivy Mike contained plutonium, californium, einsteinium, and other transuranium elements. The isotopes 249Cf, 252Cf, 253Cf, and 254Cf were observed for the first time in the debris from the explosion.[12]
Characteristics Weighable amounts of californium make it possible to determine some of its properties using macroscopic quantities. 252
Cf (2.645-year half-life) is a very strong neutron emitter and is thus extremely radioactive and harmful.[13] [14] [15] [16] [17] One microgram spontaneously emits 2,314 million neutrons per second[18] and one gram emits 39 watts of heat[19] . 249Cf is formed from the beta decay of 249Bk and most other californium isotopes are made by subjecting berkelium to intense neutron radiation in a nuclear reactor. Californium has no biological role and only a few californium compounds have been made and studied. Included among these are californium oxide (Cf2O3), californium trichloride (CfCl3) and californium oxychloride (CfOCl). The only californium ion that is stable in aqueous solution is the californium(III) cation.
Nuclear fuel cycle Actinides 244
Cm
241
Pu f
250
Cf
Halflife 243
Cmf
10–30 y
Fission products 137
Cs
90
Sr
85
Kr
Californium
5
232
U f
4n
f is for fissile
238
Pu
249
Cf f
Amf
4n
229
Cf f
246
Th
245
Cm
Cmf
250
U f
230
233
4n+1 248
U
8–24 ky
239
Pu f
32–160
231
Pa 4n+3
211–290
4n+2
Np
4n+1
99
1–2 my 247
238
Th
U
235
U f
Se
Long-lived fission products 93
135
Zr
Cs nc➔
107
Pd
80 my
Pu
79
Sn
6–23
Cmf
244
232
126
Tc
340–373
Pu
237
U
Am
242
Cm
236
234
No fission product has halflife 102 to 2×105 years
5–7 ky
243
Cm Th
Sm nc➔
431–898
251
Am
Pu
151
141–351
242
241
240
69–90 y
>7%
>5%
129
I
>1%
>.1%
fission product yield
.7–12by
Californium is produced by neutron capture on berkelium-249. Three californium isotopes with significant halflives are produced, requiring a total of 12 to 14 neutron captures on uranium-238 without nuclear fission or alpha decay. Their neutron cross sections are: Capture Th
Fission RI
Th
HL RI
(a)
250
2000
12000
251
2900
1600
4800
5500
898
252
20
44
32
1100
2.645
Cf Cf Cf
13.08
Thus 250Cf and 251Cf will be transmuted fairly quickly, with the majority fissioning at mass 251, but with a large fraction surviving to become 252Cf. The 252Cf will not be transmuted or destroyed quickly in a well-thermalized reactor, but has a short decay halflife. These isotopes decay into long-lived isotopes of curium. 252
Cf has a relatively high rate of spontaneous fission. Although still much less likely than alpha decay, this makes californium a significant neutron radiation emitter. MOX fuel containing enough curium would likely contain enough californium after use to preclude manual handling of the spent fuel or its nuclear reprocessing products with a glove box that protects against alpha and beta radiation but not against gamma radiation and especially neutron radiation.
Californium
6
Applications General 252
Cf has a number of specialized applications as a strong neutron emitter. Each milligram
of fresh californium produces 2.3 x 109 neutrons / second. Some of its uses are:[5]
[20] [21]
• neutron startup source for some nuclear reactors, calibrating instrumentation • treatment of certain cervical and brain cancers where other radiation therapy is ineffective • radiography of aircraft to detect metal fatigue • airport neutron-activation detectors of explosives • portable (non-reactor based) neutron sources for neutron activation analysis • portable metal detectors[22] • neutron moisture gauges used to find water and petroleum layers in oil wells • portable neutron source for gold and silver prospecting for on-the-spot analysis • 252Cf is used to detect ground water movement.[23] In October 2006 it was announced that three atoms ununoctium (element 118) had been identified at the Joint Institute for Nuclear Research in Dubna as the product of bombardment of californium-249 with calcium-48 ,[24] [25] [26] making this the heaviest element ever synthesized.
Military 251
Cf is famous for having a very small critical mass of 5 kg,[27] high lethality, and short period of toxic environmental irradiation relative to radioactive elements commonly used for radiation explosive weaponry, creating speculation about possible use in pocket nukes. Other weaponry uses, such as showering an area with californium, are not impossible but are seen as inhumane and are subject to inclement weather conditions and porous terrain considerations.
Literature • Guide to the Elements - Revised Edition, Albert Stwertka, (Oxford University Press; 1998) ISBN 0-19-508083-1
External links • • • •
WebElements.com - Californium [28] NuclearWeaponArchive.org - Californium [29] It's Elemental - Californium [30] Hazardous Substances Databank – Californium, Radioactive
[31]
Californium
References [1] S. G. Thompson, K. Street, Jr., A. Ghiorso, G. T. Seaborg (1950). " Element 98 (http:/ / repositories. cdlib. org/ cgi/ viewcontent. cgi?article=7072& context=lbnl)". Physical Review 78: 298. doi: 10.1103/PhysRev.78.298.2 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 78. 298. 2). . [2] S. G. Thompson, K. Street, Jr., A. Ghiorso, G. T. Seaborg (1950). " The New Element Californium (Atomic Number 98) (http:/ / www. osti. gov/ accomplishments/ documents/ fullText/ ACC0050. pdf)". Physical Review 80: 790. doi: 10.1103/PhysRev.80.790 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 80. 790). . [3] K. Street, Jr., S. G. Thompson, G. T. Seaborg (1950). " Chemical Properties of Californium (http:/ / handle. dtic. mil/ 100. 2/ ADA319899)". J. Am. Chem. Soc. 72: 4832. doi: 10.1021/ja01166a528 (http:/ / dx. doi. org/ 10. 1021/ ja01166a528). . [4] * Diamond, H. (1954). "Identification of Californium Isotopes 249, 250, 251, and 252 from Pile-Irradiated Plutonium". Phys Rev 94 (4): 1083. doi: 10.1103/PhysRev.94.1083 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 94. 1083). [5] Osborne-Lee, I.W. and Alexander, C. W. (1995). " Californium-252: A remarkable versatile radioisotope (http:/ / www. osti. gov/ bridge/ product. biblio. jsp?query_id=1& page=0& osti_id=205871)". Oak Ridge Technical Report ORNL/TM-12706. . [6] " Plutonium and Aldermaston - an historical account (http:/ / www. mod. uk/ NR/ rdonlyres/ B31B4EF0-A584-4CC6-9B14-B5E89E6848F8/ 0/ plutoniumandaldermaston. pdf)". UK Ministry of Defence. 2001-09-04. . Retrieved 2007-03-15. [7] K. Anderson, J. Pilcher, H. Wu, E. van der Bij, Z. Meggyesi, J. Adams (1999). " Neutron Irradiation Tests of an S-LINK-over-G-link System (http:/ / hep. uchicago. edu/ atlas/ tilecal/ rad/ Glink_radtest. pdf)". . [8] G. R. Burbidge et al. (1956). " PDF Californium-254 and Supernovae (http:/ / authors. library. caltech. edu/ 6553/ 1/ BURpr56. pdf)". Physical Review 103: 1145. doi: 10.1103/PhysRev.103.1145 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 103. 1145). PDF. [9] W. Baade, G. R. Burbidge, F. Hoyle, E. M. Burbidge, R. F. Christy, W. A. Fowler (1956). "Supernovae and Californium 254". Publications of the Astronomical Society of the Pacific 68: 296url = http:/ / adsabs. harvard. edu/ cgi-bin/ nph-data_query?bibcode=1956PASP. . . 68. . 296B& link_type=ARTICLE& db_key=AST& high=14365. [10] St. Temesváry (1957). "Das Element Californium-254 und die Lichtkurven der Supernovae von Typ I. Ein Beitrag zur Frage der Synthese schwerer Elemente im Kosmos". Die Naturwissenschaften 44: 321. doi: 10.1007/BF00630928 (http:/ / dx. doi. org/ 10. 1007/ BF00630928).. [11] E. Anders (1959). " Californium-254, Iron-59, and Supernovae of Type I (http:/ / adsabs. harvard. edu/ cgi-bin/ nph-data_query?bibcode=1959ApJ. . . 129. . 327A& link_type=ARTICLE& db_key=AST& high=14365)". The Astrophysical Journal 129: 327–346. doi: 10.1086/146624 (http:/ / dx. doi. org/ 10. 1086/ 146624). . [12] P. R. Fields et al. (1956). "Transplutonium Elements in Thermonuclear Test Debris". Physical Review 102 (1): 180–182. doi: 10.1103/PhysRev.102.180 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 102. 180). [13] D. A. Hicks et al. (1955). "Multiplicity of Neutrons from the Spontaneous Fission of Californium-252". Physical Review 97 (2): 564–565. doi: 10.1103/PhysRev.97.564 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 97. 564). [14] D. A. Hicks et al. (1955). "Spontaneous-Fission Neutrons of Californium-252 and Curium-244". Physical Review 98 (5): 1521–1523. doi: 10.1103/PhysRev.98.1521 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 98. 1521). [15] E. Hjalmar, H. Slätis, S.G. Thompson (1955). "Energy Spectrum of Neutrons from Spontaneous Fission of Californium-252"". Physical Review 100 (5): 1542–1543. doi: 10.1103/PhysRev.100.1542 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 100. 1542). [16] United States Patent 7118524: "Dosimetry for californium-252 (252Cf) neutron-emitting brachytherapy sources and encapsulation, storage, and clinical delivery thereof" bei www.freepatentsonline.com (http:/ / www. freepatentsonline. com/ 7118524. html). [17] Michael B. Dillon, Ronald L. Baskett, Kevin T. Foster, and Connee S. Foster (2004-03-18). " The NARAC Emergency Response Guide to Initial Airborne Hazard Estimates (https:/ / narac. llnl. gov/ uploads/ Dillon2004_NARACEmergencyResponseGuide_202990_xchnw. pdf)". National Atmospheric Release Advisory Center. . Retrieved 2008-11-14. [18] R. C. Martin, J. B. Knauer, P. A. Balo (2000, Pages 785–792). " PDF Production, Distribution, and Applications of Californium-252 Neutron Sources (http:/ / www. osti. gov/ bridge/ servlets/ purl/ 15053-AE6cnN/ native/ 15053. pdf)". Applied Radiation and Isotopes 53. doi: 10.1016/S0969-8043(00)00214-1 (http:/ / dx. doi. org/ 10. 1016/ S0969-8043(00)00214-1). PDF. [19] http:/ / www. britannica. com/ EBchecked/ topic/ 603220/ transuranium-element/ 81185/ Nuclear-properties [20] R. C. Martin and J. H. Miller. " Applications of Californium-252 Neutron Sources in Medicine, Research, and Industry (http:/ / anes. fiu. edu/ Pro/ s7Mar. pdf)" (PDF). . Retrieved 2008-11-14.
7
Californium [21] R. C. Martin (2000). " Applications and Availability of Californium-252 Neutron Sources for Waste Characterization (http:/ / www. ornl. gov/ ~webworks/ cpr/ pres/ 107270_. pdf)" (PDF). . Retrieved 2009-05-05. [22] " Will you be 'mine'? Physics key to detection (http:/ / www. pnl. gov/ news/ 2000/ 00-43. htm)". Pacific Northwest National Laboratory. 2000-10-25. . Retrieved 2007-03-21. [23] S. N. Davis et al. (2006). "Ground-Water Tracers — A Short Review". Ground Water 18 (1): 14–23. doi: 10.1111/j.1745-6584.1980.tb03366.x (http:/ / dx. doi. org/ 10. 1111/ j. 1745-6584. 1980. tb03366. x). [24] Yu. Ts. Oganessian et al. (2006). "Synthesis of the isotopes of elements 118 and 116 in the 249Cf and 245 Cm+48Ca fusion reactions". Physical Review C 74: 044602–044611. doi: 10.1103/PhysRevC.74.044602 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 74. 044602).. [25] K. Sanderson (2006-10-17). " Heaviest element made - again (http:/ / www. nature. com/ news/ 2006/ 061016/ full/ 061016-4. html)". [email protected] (Nature (journal)). . Retrieved 2006-10-19. [26] Phil Schewe and Ben Stein (2006-10-17). " Elements 116 and 118 Are Discovered (http:/ / www. aip. org/ pnu/ 2006/ 797. html)". Physics News Update. American Institute of Physics. . Retrieved 2006-10-19. [27] " Evaluation of nuclear criticality safety data and limits for actinides in transport (http:/ / europa. eu. int/ comm/ energy/ nuclear/ transport/ doc/ irsn_sect03_146. pdf)" (PDF). Institut de Radioprotection et de Sûreté Nucléaire. p. 16. . [28] [29] [30] [31]
http:/ / www. webelements. com/ webelements/ elements/ text/ Cf/ index. html http:/ / www. nuclearweaponarchive. org/ Nwfaq/ Nfaq6. html#nfaq6. 2 http:/ / education. jlab. org/ itselemental/ ele098. html http:/ / toxnet. nlm. nih. gov/ cgi-bin/ sis/ search/ r?dbs+ hsdb:@term+ @rel+ @na+ californium,radioactive
8
Article Sources and Contributors
Article Sources and Contributors Californium Source: http://en.wikipedia.org/w/index.php?oldid=304148588 Contributors: !Darkfire!6'28'14, 2D, 2help, A new name 2008, Aeroknight, Ahoerstemeier, Alansohn, AlimanRuna, Andres, Anonymous101, Antandrus, Appasionata, Atarr, Aurocker49, B07, BD2412, Beetstra, Benbest, Bender235, BillFlis, Blockhouse, BlueEarth, Bobo192, Bryan Derksen, C.A.T.S. CEO, Cadmium, Capricorn42, Carnildo, Ccmcguire, Chowbok, ChrisHodgesUK, Christacool, ChromiumCurium, Cimon Avaro, Cmprince, Conversion script, Cryptic C62, DMacks, Dan100, Darrien, David Latapie, Davidwr, Deli nk, Demiurge, DocWatson42, Duhhhhh, Duhhhhhh, Eddie tejeda, Edgar181, Emperorbma, Encyclopedia77, Epbr123, Evan Robidoux, Eyu100, Femto, Flower Priest, Flyguy649, Flying Jazz, Gaius Cornelius, Glenn4pr, Greenhousegeorge, GregorB, HazyM, Henrybauer, Hqb, Hurricane111, I80and, Infinite.08, Integrity84, J.delanoy, JNW, JWB, JWBE, Jaraalbe, Jbrennen, Jiang, Joanjoc, Josh Parris, Jrdioko, Kalamkaar, Karelj, Karlhahn, Kbh3rd, KeepItClean, Kelovy, Kingdon, Knoxville Physics, KoenigseggCCXR, Kralizec!, Kuru, Kurykh, Kwamikagami, LA2, Lightdarkness, Luna Santin, Lupin, MK8, Marc Venot, Materialscientist, Mav, Mdf, Mouser, Mumiemonstret, Myanw, Nergaal, Nick Y., Nihiltres, Nonsequiturmine, Oo64eva, Orez119, Oxymoron83, Paul August, Paulistano, PeepP, Peruvianllama, Physchim62, Piano non troppo, Polonium, Poolkris, Pras, Pt, Q43, RadiantRay, Remember, Reyk, Rjwilmsi, Roberta F., Roentgenium111, Romanm, Rwendland, Rwflammang, Saperaud, Schneelocke, Shoaler, Sin-man, Skarebo, Sl, Squids and Chips, Statue2, Stedder, Steve Crossin, Stifynsemons, Stone, Stratocracy, TEB728, TFNorman, Tagishsimon, Tcncv, Tedernst, Tenorcnj, Tetracube, That Guy, From That Show!, Thenormalyears, Tmangray, Trevor MacInnis, Uannis, UnitedStatesian, Vary, Velvetron, Versus22, Vsmith, Vuo, Walkerma, Warut, Watch37264, Weeeeeeeeeeeeee, What!?Why?Who?, Whosasking, Wiki alf, William Avery, Wknight94, Yekrats, Ytterbium2, Yyy, Zantolak, יבצ לאינד, 226 anonymous edits
Image Sources, Licenses and Contributors image:Cf-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Cf-TableImage.png License: GNU Free Documentation License Contributors: User:Ahoerstemeier File:Berkeley 60-inch cyclotron.gif Source: http://en.wikipedia.org/w/index.php?title=File:Berkeley_60-inch_cyclotron.gif License: Public Domain Contributors: Department of Energy (?) File:Elutionskurven Dy Tb Gd Eu und Cf Bk Cm Am.png Source: http://en.wikipedia.org/w/index.php?title=File:Elutionskurven_Dy_Tb_Gd_Eu_und_Cf_Bk_Cm_Am.png License: Public Domain Contributors: User:JWBE File:Kaliforn.png Source: http://en.wikipedia.org/w/index.php?title=File:Kaliforn.png License: Copyrighted free use Contributors: K. Anderson et al.
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
9
Einsteinium
1
Einsteinium californium ← einsteinium → fermium Ho ↑ Es ↓ (Uqe) Periodic Table Extended Periodic Table General Name, symbol, number Element category
einsteinium, Es, 99 actinides
Group, Period, Block
n/a, 7, f
Appearance
silver-coloured
[1]
Standard atomic weight Electron configuration Electrons per shell
−1
(252) g·mol 11
[Rn] 5f
2
7s
2, 8, 18, 32, 29, 8, 2 Physical properties
Phase
solid Density (near r.t.)
Melting point
8.84 g·cm
−3
1133 K (860 °C, 1580 °F) Atomic properties
Oxidation states Electronegativity Ionization energies
2, 3, 4 1.3 (Pauling scale) 1st: 619 kJ/mol Miscellaneous
Magnetic ordering CAS registry number
no data 7429-92-7 Most-stable isotopes
Einsteinium
2
Main article: Isotopes of einsteinium iso 252
Es
253
Es
254
Es
NA syn
syn
syn
half-life 471.7 d
20.47 d
275.7 d
DM
6.760
248
ε
1.260
252
β−
0.480
252
SF
-
39.8 d
Cf Fm
-
6.739
249
ε
0.654
254
1.090
254
6.628
250
0.288
255
α
6.436
251
SF
-
α syn
Bk
α
−
Es
DP
α
β
255
DE (MeV)
−
β
Bk Cf Fm Bk Fm Bk
-
References
Einsteinium (pronounced /aɪnˈstaɪniəm/) is a metallic synthetic element. On the periodic table, it is represented by the symbol Es and atomic number 99. It is the seventh transuranic element, and an actinoid. It was named in honor of Albert Einstein.[1]
Properties Its position on the periodic table indicates that its chemical and physical properties are similar to other metals. Though only small amounts have been made, it has been determined to be silver-colored.[1] According to tracer studies conducted at Los Alamos National Laboratory using the isotope 253Es, this element has chemical properties typical of a heavy trivalent, actinide element.[2] Like all synthetic elements, isotopes of einsteinium are extremely radioactive, and are considered highly toxic.
Production Einsteinium does not occur naturally in any measurable quantities. The modern process of creating the element starts with the irradiation of plutonium-239 in a nuclear reactor for several years. The resulting plutonium-242 isotope (in the form of the compound plutonium(IV) oxide) is mixed with aluminium and formed into pellets. The pellets are then further irradiated for approximately one year in a nuclear reactor. Another four months of irradiation is required in a different reactor. The result is a mixture of californium and einsteinium, which can then be separated.[2]
Einsteinium
Uses Aside from basic scientific research (like being a step in the production of other elements[3] ), einsteinium has no known uses.[4]
History Einsteinium was first identified in December 1952 by Albert Ghiorso along with co-workers at the University of California, Berkeley.[2] He was examining debris from the first hydrogen bomb test of November 1952 (see Operation Ivy).[1] [5] He discovered the isotope 253Es (half-life 20.5 days) that was made by the neutron capture of 15 neutrons with 238U (which then went through seven beta decays). These findings were kept secret until 1955 due to Cold War tensions.[6] [7]
Isotopes of einsteinium were produced shortly afterward at the University of California Radiation Laboratory in a nuclear fusion reaction between 14N and 238U[8] and later by intense neutron irradiation of plutonium in the Materials Testing Reactor.[9] In 1961, enough einsteinium was synthesized to prepare a microscopic amount of 253Es. This sample weighed about 0.01 mg and was measured using a special balance. The material produced was used to produce mendelevium. Further einsteinium has been produced at the Oak Ridge National Laboratory's High Flux Isotope Reactor in Tennessee by bombarding 239Pu with neutrons. Around 3 milligrams were created over a four year program of irradiation and then chemical separation from a starting 1 kg of plutonium isotope.
Isotopes Nineteen radioisotopes of einsteinium have been characterized,[10] with the most stable being 252Es with a half-life of 471.7 days. 254Es has a half-life of 275.7 days, 255Es 39.8 days and 253Es 20.47 days. All of the remaining radioactive isotopes have half-lives that are less than 40 hours, the majority of these having half-lives that are less than 30 minutes. This element also has three meta states, with the most stable being 254mEs (t½ 39.3 hours). The isotopes of einsteinium range in atomic mass from 240.069 u (240Es) to 258.100 u (258Es). The longest-lived isotope is 252Es.
Known compounds The following is a list of all known compounds of einsteinium:[11] • • • • • • •
EsBr3 einsteinium(III) bromide EsCl2 einsteinium(II) chloride EsCl3 einsteinium(III) chloride EsF3 einsteinium(III) fluoride EsI2 einsteinium(II) iodide EsI3 einsteinium(III) iodide Es2O3 einsteinium(III) oxide
3
Einsteinium
4
Literature • Stwertka, Albert (1996). A Guide to the Elements. Oxford Oxfordshire: Oxford University Press. ISBN 0195080831.
External links • • • •
Los Alamos National Laboratory - Einsteinium It's Elemental - The Element Einsteinium [13] WebElements.com - Einsteinium [14] Albert Ghiorso about the discovery [15]
[12]
References [1] Einsteinium - National Research Council Canada (http:/ / www. nrc-cnrc. gc. ca/ eng/ education/ elements/ el/ es. html). Retrieved 2 December 2007. [2] Einsteinium - Los Alamos National Laboratory (http:/ / periodic. lanl. gov/ elements/ 99. html). Retrieved 2 December 2007. [3] See Mendelevium#History [4] It's Elemental - The Element Einsteinium (http:/ / education. jlab. org/ itselemental/ ele099. html). Retrieved 2 December 2007. [5] Albert Ghiorso (2003). " Einsteinium and Fermium (http:/ / pubs. acs. org/ cen/ 80th/ einsteiniumfermium. html)". Chemical and Engineering News. . [6] Ghiorso, A. and Thompson, S. G. and Higgins, G. H. and Seaborg, G. T. and Studier, M. H. and Fields, P. R. and Fried, S. M. and Diamond, H. and Mech, J. F. and Pyle, G. L. and Huizenga, J. R. and Hirsch, A. and Manning, W. M. and Browne, C. I. and Smith, H. L. and Spence, R. W. (1955). "New Elements Einsteinium and Fermium, Atomic Numbers 99 and 100". Phys. Rev. 99: 1048–1049. doi: 10.1103/PhysRev.99.1048 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 99. 1048). [7] P. R. Fields, M. H. Studier, H. Diamond, J. F. Mech, M. G. Inghram, G. L. Pyle, C. M. Stevens, S. Fried, W. M. Manning (Argonne National Laboratory, Lemont, Illinois); A. Ghiorso, S. G. Thompson, G. H. Higgins, G. T. Seaborg (University of California, Berkeley, California): "Transplutonium Elements in Thermonuclear Test Debris", in: Physical Review 1956, 102 (1), 180–182; doi: 10.1103/PhysRev.102.180 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 102. 180). [8] Ghiorso, Albert and Rossi, G. Bernard and Harvey, Bernard G. and Thompson, Stanley G. (1954). "Reactions of U-238 with Cyclotron-Produced Nitrogen Ions". Physical Review 93: 257. doi: 10.1103/PhysRev.93.257 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 93. 257). [9] Thompson, S. G. and Ghiorso, A. and Harvey, B. G. and Choppin, G. R. (1954). "Transcurium Isotopes Produced in the Neutron Irradiation of Plutonium". Physical Review 93: 908. doi: 10.1103/PhysRev.93.908 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 93. 908). [10] Table of Isotopes decay data - LBNL Isotopes Project - LUNDS Universitet (http:/ / ie. lbl. gov/ toi/ listnuc. asp?sql=& Z=99). Retrieved 25 November 2007. [11] Chemistry : Periodic Table : einsteinium : compounds information - WebElements (http:/ / www. webelements. com/ compounds/ einsteinium). Retrieved 29 December 2008. [12] [13] [14] [15]
http:/ / periodic. lanl. gov/ elements/ 99. html http:/ / education. jlab. org/ itselemental/ ele099. html http:/ / www. webelements. com/ webelements/ elements/ text/ Es/ index. html http:/ / pubs. acs. org/ cen/ 80th/ einsteiniumfermium. html
Article Sources and Contributors
Article Sources and Contributors Einsteinium Source: http://en.wikipedia.org/w/index.php?oldid=307865124 Contributors: Addshore, AdjustShift, AdultSwim, Ahoerstemeier, Alansohn, Alba, Aleenf1, Alex Lin, AlimanRuna, AlphaEta, Bachrach44, Beetstra, Bender235, BlueDevil, Bogey97, Brian Huffman, Bryan Derksen, Bth, Caltas, Chris 73, CiTrusD, Conversion script, Cyan, DanielCD, Darrien, David Latapie, Deor, Diehard4.0, Dirac66, DocWatson42, DragonflySixtyseven, Edgar181, Egomaniac, El C, El aprendelenguas, Element16, Emperorbma, Encyclopedia77, Epbr123, Esowteric, Everyking, Eweisser, Faeriesaire, Femto, Forteblast, Fritzpoll, Func, GeorgeMoney, Gholam, Gianluigi, Giftlite, Gzornenplatz, Haham hanuka, Hak-kâ-ngìn, Hqb, Hydra Rider, II MusLiM HyBRiD II, Igoldste, Ipatrol, IvanLanin, JWBE, Jaan513, Jaked122, Jaraalbe, Jiang, Joanjoc, John, JohnCD, Jose Ramos, Jsmith86, JulieSpaulding, Kalamkaar, Kenyon, Kilo-Lima, King of Hearts, Koolade343, Kwamikagami, LA2, Logan, LonelyBeacon, Luna Santin, MK8, Marc Venot, Master runner, Materialscientist, Matty4123, Mav, Maxis ftw, Mdf, Michael Devore, Mikael V, Mike.lifeguard, Milonica, Milosci, Mormegil, Nakon, Nergaal, Nick Y., Nihiltres, NrDg, Opelio, Ossmann, Pharaoh of the Wizards, Physicistjedi, Polonium, Poolkris, Prodego, Puchiko, Rdsmith4, Reddi, Redsox345678, Reywas92, Reza kalani, Roberta F., Roentgenium111, Romanm, Rumkneebeard the Pyrate, Santăr, Saperaud, Sbharris, Schneelocke, Skunkboy74, Sl, Spiesr, Squids and Chips, Stifynsemons, Storm Rider, Stratocracy, SymlynX, Tagishsimon, Tempodivalse, ThomasK, Tide rolls, Uannis, Ucbearcatfan007, Until It Sleeps, Versus22, Violinchick1995, Vsmith, WODUP, Warut, Whosasking, Xiahou, Yamakiri, Yekrats, Yyy, Zagalejo, 249 anonymous edits
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5
Fermium
1
Fermium einsteinium ← fermium → mendelevium Er ↑ Fm ↓ (Upn) Periodic Table Extended Periodic Table General Name, symbol, number Element category
fermium, Fm, 100 actinides
Group, Period, Block
n/a, 7, f
Appearance
unknown, probably silvery white or metallic gray Standard atomic weight Electron configuration Electrons per shell
−1
(257) g·mol 12
[Rn] 5f
2
7s
2, 8, 18, 32, 30, 8, 2 Physical properties
Phase
solid
Melting point
1800 K (1527 °C, 2781 °F) Atomic properties
Oxidation states Electronegativity Ionization energies
2, 3 1.3 (Pauling scale) 1st: 627 kJ/mol Miscellaneous
CAS registry number
7440-72-4 Most-stable isotopes
Fermium
2
Main article: Isotopes of fermium iso 252
Fm
253
Fm
255
Fm
257
Fm
NA syn
syn
syn
syn
half-life 25.39 h
3d
20.07 h
100.5 d
DM
DE (MeV)
DP
SF
-
α
7.153
248
ε
0.333
253
α
7.197
249
SF
-
α
7.241
251
6.864
253
α SF
-
Cf Es Cf
Cf Cf
-
References
Fermium (pronounced /ˈfɜrmiəm/) is a synthetic element with the symbol Fm and atomic number 100. A highly radioactive metallic transuranic element of the actinide series, fermium is made by bombarding plutonium with neutrons and is named after nuclear physicist Enrico Fermi. Fermium is the eighth transuranic element.
Characteristics Only small amounts of fermium have ever been produced or isolated. Thus relatively little is known about its chemical properties. Only the (III) oxidation state of the element appears to exist in aqueous solution. 254Fm and heavier isotopes can be synthesized by intense neutron bombardment of lighter elements (especially uranium and plutonium). During this, successive neutron captures mixed with beta decays build the fermium isotope. The intense neutron bombardment conditions needed to create fermium exist in thermonuclear explosions and can be replicated in the laboratory (such as in the High Flux Isotope Reactor at Oak Ridge National Laboratory). The Electron shell diagram of fermium synthesis of element 102 (nobelium) was confirmed 250 when Fm was chemically identified. Like all synthetic elements is it extremely radioactive and highly toxic.
Fermium
Uses There are no known uses of fermium outside of basic research.
History Fermium (after Enrico Fermi) was first discovered by a team led by Albert Ghiorso in 1952. The team found 255Fm in the debris of the first hydrogen bomb explosion (see Operation Ivy). That isotope was created when 238U combined with 17 neutrons in the intense temperature and pressure of the explosion (eight beta decays also occurred to create the element). The work was overseen by the University of California Radiation Laboratory, Argonne National Laboratory, and Los Alamos Scientific Laboratory. All these findings were kept secret until 1955 due to Cold War tensions.[1] Samples of sea coral impacted from the first thermonuclear explosion of November 1952 were used.[2] In late 1953 and early 1954 a team from the Nobel Institute of Physics in Stockholm bombarded a 238U target with 16O ions, producing an alpha-emitter with an atomic weight of ~250 and with 100 protons (in other words, element 250100).[3] The Nobel team did not claim discovery until 1954. The isotope they produced was later positively identified as 250 Fm.
Isotopes 17 radioisotopes of fermium have been characterized, with the most stable being 257Fm with a half-life of 100.5 days, 253Fm with a half-life of 3 days, 252Fm with a half-life of 25.39 hours, and 255Fm with a half-life of 20.07 hours. All of the remaining radioactive isotopes have half-lives that are less than 5.4 hours, and the majority of these have half-lives that are less than 3 minutes. This element also has 1 meta state, 250mFm (t½ 1.8 seconds). The isotopes of fermium range in atomic weight from 242.073 u (242Fm) to 259.101 u (259Fm).
References [1] Ghiorso, A.; Thompson, S. G.; Higgins, G. H. ; Seaborg, G. T.; Studier, M. H.; Fields, P. R.; Fried, S. M.; Diamond, H.; Mech, J. F.; Pyle, G. L.; Huizenga, J. R.; Hirsch, A.; Manning, W. M.; Browne, C. I.; Smith, H. L.; Spence, R. W. (1955). " New Elements Einsteinium and Fermium, Atomic Numbers 99 and 100 (http:/ / prola. aps. org/ abstract/ PR/ v99/ i3/ p1048_1)". Physical Review 99: 1048–1049. doi: 10.1103/PhysRev.99.1048 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 99. 1048). . [2] Albert Ghiorso (2003). " Einsteinium and Fermium (http:/ / pubs. acs. org/ cen/ 80th/ einsteiniumfermium. html)". Chemical and Engineering News. . [3] Atterling, Hugo; Forsling, Wilhelm; Holm, Lennart W.; Melander, Lars; Åström, Björn (1954). "Element 100 Produced by Means of Cyclotron-Accelerated Oxygen Ions". Physical Review 95: 585–586. doi: 10.1103/PhysRev.95.585.2 (http:/ / dx. doi. org/ 10. 1103/ PhysRev. 95. 585. 2).
• Los Alamos National Laboratory - Fermium (http:/ / periodic. lanl. gov/ elements/ 100. html) • Guide to the Elements - Revised Edition, Albert Stwertka, (Oxford University Press; 1998) ISBN 0-19-508083-1 • It's Elemental - Fermium (http:/ / education. jlab. org/ itselemental/ ele100. html)
3
Fermium
External links • WebElements.com - Fermium (http:/ / www. webelements. com/ webelements/ elements/ text/ Fm/ index. html)
4
Article Sources and Contributors
Article Sources and Contributors Fermium Source: http://en.wikipedia.org/w/index.php?oldid=307492879 Contributors: Achaemenes, Ahoerstemeier, Ale jrb, AlimanRuna, Arkuat, Beetstra, Bentley4, Bkell, BlueEarth, Brian Huffman, Bryan Derksen, Carnildo, ChemNerd, Chris Roy, Conversion script, Darrien, David Latapie, Dina, Edgar181, Egomaniac, Emperorbma, Femto, Glenn4pr, Hak-kâ-ngìn, Hqb, Ideyal, Jaan513, Jaraalbe, Jiang, John, Justfred, Karelj, Kelovy, Kjramesh, Kwamikagami, Lainex, Looxix, Mav, Mdf, Mimihitam, Nergaal, Nick Y., PierceG, Polonium, Poolkris, Pras, Remember, Rjwilmsi, Roberta F., Sandman, Saperaud, Schneelocke, Sl, Stifynsemons, Stratocracy, Tagishsimon, Uannis, Vsmith, Whosasking, Willking1979, Yekrats, Yyy, 63 anonymous edits
Image Sources, Licenses and Contributors image:Fm-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Fm-TableImage.png License: GNU Free Documentation License Contributors: User:Ahoerstemeier Image:Electron shell 100 Fermium.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_100_Fermium.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80
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5
Mendelevium
1
Mendelevium fermium ← mendelevium → nobelium Tm ↑ Md ↓ (Upu) Periodic Table Extended Periodic Table General Name, symbol, number Element category
mendelevium, Md, 101 actinides
Group, Period, Block
n/a, 7, f
Appearance
unknown, probably silvery white or metallic gray Standard atomic weight Electron configuration Electrons per shell
−1
(258) g·mol 13
[Rn] 5f
2
7s
2, 8, 18, 32, 31, 8, 2 Physical properties
Phase
solid
Melting point
1100 K (827 °C, 1521 °F) Atomic properties
Oxidation states Electronegativity Ionization energies
2, 3 1.3 (Pauling scale) 1st: 635 kJ/mol Miscellaneous
Magnetic ordering CAS registry number
no data 7440-11-1 Most-stable isotopes
Mendelevium
2
Main article: Isotopes of mendelevium iso 257
Md
NA syn
half-life 5.52 h
DM
DE (MeV)
DP
ε
0.406
257
α
7.558
253
SF
-
Fm Es
-
258
syn
51.5 d
ε
1.230
260
syn
31.8 d
SF
-
α
7.000
256
ε
-
260
1.000
260
Md Md
−
β
258
Fm
Es Fm No
References
Mendelevium (pronounced /ˌmɛndəˈlɛviəm/ or /ˌmɛndəˈliːviəm/) is a synthetic element with the symbol Md (formerly Mv) and the atomic number 101. A metallic radioactive transuranic element of the actinides, mendelevium is synthesized by bombarding einsteinium with alpha particles and was named after Dmitri Mendeleev, who was responsible for the Periodic Table.
Characteristics Researchers have shown that mendelevium has a moderately stable dipositive (II) oxidation state in addition to the more characteristic (for actinide elements) tripositive (III) oxidation state, the latter being the more dominantly exhibited state in an aqueous solution (Chromatography being the process used). Sometimes, Mendelevium can even be shown to exhibit a monopositive (I) state. [1] 256Md has been used to find out some of the chemical properties of this element while in an aqueous solution. There are no other known uses of mendelevium and only trace amounts of the element have ever been produced. Other isotopes of Mendelevium, all radioactive have been discovered, with 258Md being the most stable with a two-month half-life (about 55 days). [2] Other isotopes range from 248 to 258 mass numbers and half-lives from a few seconds to about 55 days. The original 256Md had a half-life of 76 minutes. [3] It is assumed that Mendelevium's standard state is solid at 298K and its classification is metallic. [4]
History Mendelevium (for Dmitri Mendeleev, surname commonly transliterated into Latin script as Mendeleev, Mendeleyev, Mendeléef, or even Mendelejeff, and first name sometimes transliterated as Dmitry or Dmitriy) was first synthesized by Albert Ghiorso (team leader), Glenn T. Seaborg, Gregory R. Choppin, Bernard G. Harvey, and Stanley G. Thompson in early 1955 at the University of California, Berkeley. The team produced 256Md (half-life of 76 minutes) when they bombarded an 253Es target with alpha particles (helium nuclei) in the Berkeley Radiation Laboratory's 60-inch cyclotron (256Md was the first element to be synthesized one-atom-at-a-time). Element 101 was the ninth transuranic element
Mendelevium
3
synthesized. The first 17 atoms of this element were created and analyzed using the ion-exchange adsorption-elution method. During the process, Mendelevium behaved very much like Thulium, its naturally-occurring homologue. [5]
Isotopes 15 radioisotopes of mendelevium have been characterized, with the most stable being 258 Md with a half-life of 51.5 days, 260Md with a half-life of 31.8 days, and 257Md with a half-life of 5.52 hours. All of the remaining radioactive isotopes have half-lives that are less than 97 minutes, and the majority of these have half-lives that are less than 5 minutes. This element also has 1 meta state, 258mMd (t½ 57 minutes). The isotopes of mendelevium range in atomic weight from 245.091 u (245Md) to 260.104 u (260Md).
References • Los Alamos National Laboratory - Mendelevium
[6]
• Guide to the Elements - Revised Edition, Albert Stwertka, (Oxford University Press; 1998) ISBN 0-19-508083-1 • It's Elemental - Mendelevium [7] [1] [2] [3] [4] [5] [6] [7]
http:/ / www. answers. com/ topic/ mendelevium http:/ / library. eb. co. uk/ eb/ article-9051976 http:/ / www. answers. com/ topic/ mendelevium http:/ / www. webelements. com/ mendelevium/ http:/ / library. eb. co. uk/ eb/ article-9051976 http:/ / periodic. lanl. gov/ elements/ 101. html http:/ / education. jlab. org/ itselemental/ ele101. html
External links • WebElements.com - Mendelevium (http:/ / www. webelements. com/ webelements/ elements/ text/ Md/ index. html) • Environmental Chemistry- Md info (http:/ / environmentalchemistry. com/ yogi/ periodic/ Md. html)
Article Sources and Contributors
Article Sources and Contributors Mendelevium Source: http://en.wikipedia.org/w/index.php?oldid=302987020 Contributors: 1exec1, Ahoerstemeier, AlimanRuna, Arkuat, Beetstra, Benbest, Bill, Bonzostar, Brian Huffman, Bryan Derksen, Carnildo, ClanCC, Conversion script, Darrien, David Latapie, Edgar181, El C, Emperorbma, Femto, Gdommett, Glenn4pr, Hak-kâ-ngìn, Hede2000, Hqb, Ideyal, Infinite.08, Jaan513, Jaraalbe, Jiang, Joanjoc, John beaker, Karl-Henner, Kelovy, Kwamikagami, LittleOldMe, Malcolm Farmer, Mark Musante, Mav, Mdf, Michel32Nl, Mimihitam, Narval, Nergaal, Nick Y., Nihiltres, OverMyHead, PeepP, Polonium, Poolkris, Pras, Pstanton, Remember, Reza kalani, Roberta F., Roentgenium111, Romanm, Saperaud, Schneelocke, SimonP, Sl, Squids and Chips, Stifynsemons, Stratocracy, Tagishsimon, Tasc, ThaddeusB, Uannis, VASANTH S.N., Vcelloho, Videogamer344, Vsmith, Wiwaxia, Writtenright, Yekrats, 81 anonymous edits
Image Sources, Licenses and Contributors image:Md-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Md-TableImage.png License: GNU Free Documentation License Contributors: User:Ahoerstemeier
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4
Nobelium
1
Nobelium mendelevium ← nobelium → lawrencium Yb ↑ No ↓ (Upb) Periodic Table Extended Periodic Table General Name, symbol, number Element category Group, Period, Block Appearance
nobelium, No, 102 actinides n/a, 7, f unknown, probably metallic
Standard atomic weight Electron configuration Electrons per shell
−1
[259] g·mol 14
[Rn] 5f
2
7s
2, 8, 18, 32, 32, 8, 2 Physical properties
Phase
solid Atomic properties
Oxidation states Electronegativity Ionization energies
2, 3 (Pauling scale) 1st: 641.6 kJ/mol 2nd: 1254.3 kJ/mol 3rd: 2605.1 kJ/mol Miscellaneous
CAS registry number
10028-14-5 Most-stable isotopes
Nobelium
2
Main article: Isotopes of nobelium iso
NA
half-life
262
syn
5 ms
260
syn
106 ms
259
syn
58m
No No No
DM
DE (MeV)
DP
SF SF 75% α
7.69,7.61,7.53....
25% ε
255
Fm
259
Md
258
syn
1.2 ms
SF
257
syn
25 s
α
8.32,8.22
253
256
syn
2.91 s
99.5% α
8.45,8.40
252
61% α
8.12,8.08,7.93
251
39% ε
2.012
255
No No No
Fm Fm
0.5% f 255
No
syn
3.1 m
Fm Md
254
m2
syn
198 µs
γ
254
m1
254
m1
syn
275 ms
γ
250
g
254
g
syn
51 s
253
m
syn
43.5 µs
γ
253
g
syn
1.62 m
α
8.14,8.06,8.04,8.01
249
75% α
8.42,8.37
248
8.62,8.58
247
No No No No
253
No
252
m
syn
110 ms
252
g
syn
2.44 s
No No
No No
No
Fm
Fm
25% SF 251
syn
0.76 s
α
250
syn
43 µs
SF
250
syn
3.7 µs
SF
No Nom Nog
Fm
References
Nobelium (pronounced /noʊˈbɛliəm/ or /noʊˈbiːliəm/) is a synthetic element with the symbol No and atomic number 102. It was first correctly identified in 1956 by scientists at the Flerov Laboratory of Nuclear Reactions in Dubna, Russia. Little is known about the element but limited chemical experiments have shown that it forms a stable divalent ion in solution as well as the predicted trivalent ion that is associated with its presence as one of the actinoids.
Nobelium
3
Discovery profile The discovery of element 102 was first announced by physicists at the Nobel Institute in Sweden in 1957. The team reported that they created an isotope with a half-life of 10 minutes, decaying by emission of an 8.5 MeV alpha particle, after bombarding 244Cm with 13 C nuclei. The activity was assigned to 251102 or 253102. The scientists proposed the name nobelium (No) for the new element. Later they retracted their claim and associated the activity to background effects. The synthesis of element 102 was then claimed in April 1958 at the University of California, Berkeley by Albert Ghiorso, Glenn T. Seaborg, John R. Walton and Torbjørn Sikkeland. The team used the new heavy-ion linear accelerator (HILAC) to bombard a curium target (95% 244 Cm and 5% 246Cm) with 13C and 12C ions. They were unable to confirm the 8.5 MeV activity claimed by the Swedes but were instead able to detect decays from 250Fm, supposedly the daughter of 254102, which had an apparent half-life of ~3 s. In 1959 the team continued their studies and claimed that they were able to produce an isotope that decayed predominantly by emission of an 8.3 MeV alpha particle, with a half-life of 3 s with an associated 30% spontaneous fission branch. The activity was initially assigned to 254No but later changed to 252No. The Berkeley team decided to adopt the name nobelium for the element. 24496Cm
+
126C
→
*
256102No
→
252102No
+4
10n
Further work in 1961 on the attempted synthesis of element 103 (see lawrencium) produced evidence for a Z=102 alpha activity decaying by emission of an 8.2 MeV particle with a half-life of 15 s, and assigned to 255No. Following initial work between 1958-1964, in 1966, a team at the Flerov Laboratory of Nuclear Reactions (FLNR) reported that they had been able to detect 250Fm from the decay of a parent nucleus (254No) with a half-life of ~50s, in contradiction to the Berkeley claim. Furthermore, they were able to show that the parent decayed by emission of 8.1 MeV alpha particles with a half-life of ~35 s. 23892U
+
2210Ne
→
*
260102No
→
254102No
+6
10n
In 1969, the Dubna team carried out chemical experiments on element 102 and concluded that it behaved as the heavier homologue of Ytterbium. The Russian scientists proposed the name joliotium (Jo) for the new element. Later work in 1967 at Berkeley and 1971 at Oak Ridge fully confirmed the discovery of element 102 and clarified earlier observations. In 1992, the IUPAC-IUPAP Transfermium Working Group (TWG) discovery and concluded that only the Dubna work from 1966 assigned decays to Z=102 nuclei at the time. The Dubna team recognised as the discoverers of nobelium although it is possible Berkeley in 1959.
assessed the claims of correctly detected and are therefore officially that it was detected at
Nobelium
4
Naming Element 102 was first named nobelium (No) by its claimed discoverers in 1957 by scientists at the Nobel Institute in Sweden. The name was later adopted by Berkeley scientists who claimed its discovery in 1959. The International Union of Pure and Applied Chemistry (IUPAC) officially recognised the name nobelium following the Berkeley results. In 1994, and subsequently in 1997, the IUPAC ratified the name nobelium (No) for the element on the basis that it had become entrenched in the literature over the course of 30 years and that Alfred Nobel should be commemorated in this fashion.
Electronic structure Nobelium is element 102 in the Periodic Table. The two forms of the projected electronic structure are: Bohr model Quantum mechanical model
2, 8, 18, 32, 32, 8, 2 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f14
Physical properties The appearance of this element is unknown, however it is most likely silvery-white or gray and metallic. If sufficient amounts of nobelium were produced, it would pose a radiation hazard. Some sources quote a melting point of 827°C for nobelium but this cannot be substantiated from an official source and seems implausible regarding the requirements of such a measurement. However, the 1st, 2nd and 3rd ionization energies have been measured. In addition, an electronegativity value of 1.3 is also sometimes quoted. This is most definitely only an estimate since a true value can only be determined using a chemical compound of the element and no such compounds exist for nobelium.
Experimental chemistry Aqueous phase chemistry First experiments involving nobelium assumed that it predominantly formed a +III state like earlier actinoids. However, it was later found that nobelium forms a stable +II state in solution, although it can be oxidised to an oxidising +III state.[1] A reduction potential of -1.78 V has been measured for the No3+ ion. The hexaaquanobelium(II) ion has been determined to have an ionic radius of 110 pm.
Nobelium
5
Summary of compounds and (complex) ions Formula
Names(s)
[No(H2O)6]3+
hexaaquanobelium(III)
[No(H2O)6]2+
hexaaquanobelium(II)
Isotopes Seventeen radioisotopes of nobelium have been characterized, with the most stable being 259 No with a half-life of 58 minutes. Longer half-lives are expected for the as-yet-unknown 261 No and 263No. An isomeric level has been found in 253No and K-isomers have been found in 250No, 252No and 254No to date.
Synthesis of isotopes as decay products Isotopes of nobelium have also been identified in the decay of heavier elements. Observations to date are summarised in the table below: Evaporation Residue
Observed No isotope
262
262
Lr
No
269
265
261
257
267
263
259
255
Hs, Hs,
Sg, Sg,
Rf
No
Rf
No
254
254
Lr
No
261
257
264
260
Sg, Hs,
253
Rf Sg,
No
256
252
Rf
No
255
251
Rf
No
Chronology of isotope discovery Isotope
Year discovered
Discovery reaction
250
2001
204
250
2006
204
251
1967
244
252
1959
244
252
~2002
206
253
1967
242
253
1971
249
254
1966
243
254
1967?
246
Nom Nog No Nog Nom Nog Nom Nog Nom1
Pb(48Ca,2n) Pb(48Ca,2n) Cm(12C,5n) Cm(12C,4n) Pb(48Ca,2n) Pu(16O,5n),239Pu(18O,4n) [2]
Cf(12C,4n)
Am(15N,4n) Cm(13C,5n),246Cm(12C,4n)
Nobelium
6
254
~2003
208
255
1967
246
256
1967
248
257
1961? , 1967
248
258
1967
248
259
1973
248
260
?
254
22
261
unknown
262
1988
254
22
Nom2 No No No No No No No No
Pb(48Ca,2n) Cm(13C,4n),248Cm(12C,5n) Cm(12C,4n),248Cm(13C,5n) Cm(13C,4n) Cm(13C,3n) Cm(18O,α3n) Es +
Es +
Ne,18O,13C - transfer
Ne - transfer (EC of
262
Lr)
Isomerism in nobelium nuclides 254
No The study of K-isomerism was recently studied by physicists at the University of Jyväskylä physics laboratory (JYFL). They were able to confirm a previously reported K-isomer and detected a second K-isomer. They assigned spins and parities of 8- and 16+ to the two K-isomers. 253
No In 1971, Bemis et al. was able to determine an isomeric level decaying with a half-life of 31 µs from the decay of 257Rf. This was confirmed in 2003 at the GSI by also studying the decay of 257Rf. Further support in the same year from the FLNR appeared with a slightly higher half-life of 43.5 µs, decaying by M2 gamma emission to the ground state. 252
No In a recent study by the GSI into K-isomerism in even-even isotopes, a K-isomer with a half-life of 110 ms was detected for 252No. A spin and parity of 8- was assigned to the isomer. 250
No In 2003, scientists at the FLNR reported that they had been able to synthesise 249No which decayed by SF with a half-life of 54µs. Further work in 2006 by scientists at the ANL showed that the activity was actually due to a K-isomer in 250No. The ground state isomer was also detected with a very short half-life of 3.7µs.
Chemical yields of isotopes Cold fusion The table below provides cross-sections and excitation energies for cold fusion reactions producing nobelium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
CN
1n
2n
48
208
256
254
48
207
255
253
48
206
254
Ca
Ca
Ca
Pb
Pb
Pb
4n
No: 2050 nb ; 22.3 MeV
No
No: 1310 nb ; 22.4 MeV
No
No
3n
253
No: 58 nb ; 23.6 MeV
252
No: 515 nb ; 23.3 MeV
251
No: 30 nb ; 30.7 MeV
250
No: 260 pb ; 43.9 MeV
Nobelium
48
Ca
7
204
Pb
250
252
No:13.2 nb ; 23.2 MeV
No
Hot fusion The table below provides cross-sections and excitation energies for hot fusion reactions producing nobelium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
CN
3n
4n
5n
6n
26
232
258
254
253
252
22
238
260
256
255
254
22
236
258
254
253
252
Mg Ne Ne
Th U U
No
No:1.6 nb
No
No:40 nb
No
No:7 nb
No:9 nb No:200 nb No:25 nb
No:8 nb No:15 nb No:15 nb
Retracted isotopes In 2003, scientists at the FLNR claimed to have discovered the lightest known isotope of nobelium. However, subsequent work showed that the 54 µs activity was actually due to 250 No and the isotope 249No was retracted.
Notes • Los Alamos National Laboratory - Nobelium
[3]
• Guide to the Elements - Revised Edition, Albert Stwertka, (Oxford University Press; 1998) ISBN 0-19-508083-1 • It's Elemental - Nobelium [4]
External links • WebElements.com - Nobelium
[5]
References [1] Toyoshima, Atsushi (2009). "Oxidation of Element 102, Nobelium, with Flow Electrolytic Column Chromatography on an Atom-at-a-Time Scale". Journal of the American Chemical Society: 090610145759060. doi: 10.1021/ja9030038 (http:/ / dx. doi. org/ 10. 1021/ ja9030038). edit (http:/ / en. wikipedia. org/ wiki/ Template:cite_doi/ 10. 1021. 2fja9030038) [2] [3] [4] [5]
see rutherfordium http:/ / periodic. lanl. gov/ elements/ 102. html http:/ / education. jlab. org/ itselemental/ ele102. html http:/ / www. webelements. com/ webelements/ elements/ text/ No/ index. html
Article Sources and Contributors
Article Sources and Contributors Nobelium Source: http://en.wikipedia.org/w/index.php?oldid=308809697 Contributors: Ahoerstemeier, Alfio, AlimanRuna, Andres, Atjesse, Bender235, Black-Velvet, Bobo192, BostonMA, Brian Huffman, Bryan Derksen, CO, Carnildo, ChemNerd, Conversion script, Darrien, David Latapie, DesertAngel, Discospinster, Drjezza, Edgar181, Emperorbma, Everyking, Fangfufu, Femto, Glenn4pr, Hak-kâ-ngìn, Hede2000, IRP, Ideyal, J.delanoy, J04n, Jaraalbe, Jiang, Joanjoc, John, Karelj, Karl-Henner, Kwamikagami, LA2, LeaveSleaves, LizardWizard, Marc9510000, Mav, Mdf, Mimihitam, Miquonranger03, Miss Madeline, Mmxx, Montrealais, Nergaal, Nihiltres, Polonium, Poolkris, Pras, PrimeCupEevee, Remember, Renrenren, Reza kalani, Roberta F., Roentgenium111, Rursus, Santăr, Saperaud, Schneelocke, Seegoon, Skater, Sl, Snigbrook, Stifynsemons, Swalot, Tagishsimon, Tetracube, Tom harrison, Uannis, Undead warrior, VASANTH S.N., Van helsing, Vsmith, Vuo, Warut, Watch37264, Yekrats, Yyy, 90 anonymous edits
Image Sources, Licenses and Contributors image:No-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:No-TableImage.png License: GNU Free Documentation License Contributors: User:Ahoerstemeier
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
8
Lawrencium
1
Lawrencium nobelium ← lawrencium → rutherfordium Lu ↑ Lr ↓ (Upt) Periodic Table Extended Periodic Table General Name, symbol, number Element category
lawrencium, Lr, 103 transition metals
Group, Period, Block
n/a, 7, d
Appearance
unknown, probably silvery white or metallic gray Standard atomic weight Electron configuration Electrons per shell
−1
[262] g·mol 14
[Rn] 5f
2
1
7s 7p
2, 8, 18, 32, 32, 8, 3 Physical properties
Phase
presumably a solid Atomic properties
Oxidation states
3
Ionization energies
1st: 443.8 kJ/mol 2nd: 1428.0 kJ/mol 3rd: 2219.1 kJ/mol Miscellaneous
CAS registry number
22537-19-5 Most-stable isotopes
Lawrencium
2
Main article: Isotopes of lawrencium iso
NA
half-life
DM
DE (MeV)
DP
262
syn
3.6 h
EC
261
syn
44 m
SF/EC?
260
syn
2.7 m
alpha
8.04
256
259
syn
6.2 s
78% alpha
8.44
255
Lr Lr Lr Lr
262
No
Md Md
22% SF 258
syn
4.1 s
alpha
8.68,8.65,8.62,8.59
254
257
syn
0.65 s
alpha
8.86,8.80
253
256
syn
27 s
alpha
8.62,8.52,8.32...
252
255
syn
21.5 s
alpha
8.43,8.37
251
254
syn
13 s
78% alpha
8.46,8.41
250
Lr Lr Lr Lr Lr
22% EC
Md Md Md Md Md
254
No
253
m
syn
0.57 s
alpha
8.79
249
253
g
syn
1.49 s
92% alpha
8.72
249
9.02,8.97
248
Lr Lr
Md Md
8% SF 252
Lr
syn
0.36 s
alpha
Md
References
Lawrencium (pronounced /ləˈrɛnsiəm/) is a radioactive synthetic element with the symbol Lr (formerly Lw) and atomic number 103. Its most stable known isotope is 262Lr, with a half-life of approximately 3.6 hours. Little is known of the chemistry but there is strong evidence for the formation of a trivalent ion in aqueous solution, confirming lawrencium's place as the last member of the actinoids. Although lawrencium is often placed as the last member of the 5f-block, it can also be regarded as the first member of the 6d-block (see extended periodic table).
Official discovery Lawrencium was made by Albert Ghiorso, Torbjørn Sikkeland, Almon Larsh, and Robert M. Latimer on February 14, 1961 at the Lawrence Radiation Laboratory (now called Lawrence Berkeley National Laboratory) on the University of California, Berkeley campus. It was produced by bombarding a three milligram target composed of three isotopes of californium with boron-10 and B-11 ions in the Heavy Ion Linear Accelerator (HILAC). The Berkeley team reported that the isotope 257103 was detected in this manner and decayed by emitting an 8.6 MeV alpha particle with a half-life of ~8 seconds. The assignment was later corrected to 258Lr. 25298Cf
+
115B
→
263−x103Lr
→
258103Lr
+5
10n
Lawrencium
3
The team suggested the name lawrencium (Lw) for the new element. In 1967, researchers in Dubna, Russia reported that they were not able to confirm an alpha emitter with a half-life of 8 seconds as 257103. This assignment has since been changed to 258 Lr. Instead, they reported a 45s activity assigned to 256Lr. 24395Am
+
188O
→
261−x103Lr
→
256103Lr
+5
10n
Further work in 1969 indicated an actinoid chemistry for the new element founded by Travis Anselm in 8B. In 1971, the team at the University of California performed a whole series of experiments aimed at measuring the decay properties of lawrencium isotopes with mass numbers from 255-260. In 1992, The IUPAC/IUPAP Transfermium Working Group (TWG) officially recognised the Dubna and Berkeley teams as co-discoverers of lawrencium.
Naming The origin of the name, preferred by the American Chemical Society, is in reference to Ernest O. Lawrence, inventor of the cyclotron. The symbol Lw originally was used but in 1963 it was changed to Lr. In August 1997 the International Union of Pure and Applied Chemistry (IUPAC) ratified the name lawrencium and symbol Lr during a meeting in Geneva. Lawrencium has also been referred to as eka-lutetium.
Electronic structure Lawrencium is element 103 in the Periodic Table. The two forms of the projected electronic structure are: Bohr model Quantum mechanical model
2, 8, 18, 32, 32, 9, 2 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f146d1
There has been a suggestion that the electron configuration could be [Rn]7s25f147p1 instead. Direct measurement is impossible, however, and calculations have given conflicting results.[1]
Physical characteristics The appearance of this element is unknown, however it is most likely silvery-white or gray and metallic. If sufficient amounts of lawrencium were produced, it would pose a radiation hazard. Contrary to some sources, bulk properties of this element, such as the melting point, have not been possible to measure to date. However, the 1st, 2nd and 3rd ionization energies have been measured.
Lawrencium
4
Periodic classification A strict correlation between periodic table blocks and electron configuration for neutral atoms would describe lawrencium as a transition metal because it should be classed as a d-block element. However, it is classified as an actinoid according to IUPAC recommendations.[2]
Experimental chemistry Gas phase chemistry The first gas phase studies were reported in 1969 by a team at the Flerov Laboratory of Nuclear Reactions (FLNR). They used the reaction 243Am+18O to produce lawrencium nuclei which reacted with a stream of chlorine gas to form a volatile chloride component. The product was assigned to 256LrCl3 and confirmed that lawrencium was a typical actinide.
Aqueous phase chemistry The first liquid phase studies were reported in 1970 by the team at the LBNL. They used the reaction 249Cf+11B to produce lawrencium nuclei. They were able to show that lawrencium formed a trivalent ion, similar to other actinides but in stark contrast to nobelium. Further work in 1988 confirmed the formation of a trivalent lawrencium(III) ion using anion-exchange chromatography using α-hydroxyisobutyrate (α-HIB) complex. Comparison of the elution time with other actinides allowed a determination of 88.6 pm for the ionic radius for Lr3+. Attempts to reduce Lr(III) to Lr(I) using the potent reducing agent hydroxylamine hydrochloride were unsuccessful.
Summary of compounds and complex ions Formula LrCl3
Names(s) lawrencium trichloride ; lawrencium(III) chloride
Isotopes Twelve isotopes of lawrencium have been synthesized with 262Lr being the longest-lived and heaviest, with a half-life of 216 minutes. 252Lr is the lightest isotope known to date.
Lawrencium
History of synthesis of isotopes by cold fusion 205
Tl(50Ti,xn)255-xLr (x=2?)
This reaction was studied in a series of experiments in 1976 by Yuri Oganessian and his team at the FLNR. Evidence was provided for the formation of 253Lr in the 2n exit channel. 203
Tl(50Ti,xn)253-xLr
This reaction was studied in a series of experiments in 1976 by Yuri Oganessian and his team at the FLNR. 208
Pb(48Ti,pxn)255-xLr (x=1?)
This reaction was reported in 1984 by Yuri Oganessian at the FLNR. The team was able to detect decays of 246Cf, a descendant of 254Lr. 208
Pb(45Sc,xn)253-xLr
This reaction was studied in a series of experiments in 1976 by Yuri Oganessian and his team at the FLNR. Results are not readily available. 209
Bi(48Ca,xn)257-xLr (x=2)
This reaction has been used to study the spectroscopic properties of 255Lr. The team at GANIL used the reaction in 2003 and the team at the FLNR used it between 2004-2006 to provide further information for the decay scheme of 255Lr. The work provided evidence for an isomeric level in 255Lr.
History of synthesis of isotopes by hot fusion 243
Am(18O,xn)261-xLr (x=5)
This reaction was first studied in 1965 by the team at the FLNR. They were able to detect a 45s activity assigned to 256Lr or 257Lr. Later work suggests an assignment to 256Lr. Further studies in 1968 produced an 8.35-8.60 MeV alpha activity with a half-life of 35s. This activity was also initially assigned to 256Lr or 257Lr and later to solely 256Lr. 243
Am(16O,xn)259-xLr (x=4)
This reaction was studied in 1970 by the team at the FLNR. They were able to detect an 8.38 MeV alpha activity with a half-life of 20s. This was assigned to 255Lr.
5
Lawrencium 248
Cm(15N,xn)263-xLr (x=3,4,5)
This reaction was studied in 1971 by the team at the LBNL in their large study of lawrencium isotopes. They were able to assign alpha activities to 260Lr,259Lr and 258Lr from the 3-5n exit channels. 248
Cm(18O,pxn)265-xLr (x=3,4)
This reaction was studied in 1988 at the LBNL in order to assess the possibility of producing 262Lr and 261Lr without using the exotic 254Es target. It was also used to attempt to measure an EC branch in 261mRf from the 5n exit channel. After extraction of the Lr(III) component, they were able to measure the spontaneous fission of 261Lr with an improved half-life of 44 minutes. The production cross-section was 700 pb. On this basis, a 14% EC branch was calculated if this isotope was produced via the 5n channel rather than the p4n channel. A lower bombarding energy (93 MeV c.f. 97 MeV) was then used to measure the production of 262Lr in the p3n channel. The isotope was successfully detected and a yield of 240 pb was measured. The yield was lower than expected compared to the p4n channel. However, the results were judged to indicate that the 261Lr was most likely produced by a p3n channel and an upper limit of 14% for the EC branch of 261mRf was therefore suggested. 246
Cm(14N,xn)260-xLr (x=3?)
This reaction was studied briefly in 1958 at the LBNL using an enriched 244Cm target (5% 246 Cm). They observed a ~9 MeV alpha activity with a half-life of ~0.25 seconds. Later results suggest a tentative assignment to 257Lr from the 3n channel 244
Cm(14N,xn)258-xLr
This reaction was studied briefly in 1958 at the LBNL using an enriched 244Cm target (5% 246 Cm). They observed a ~9 MeV alpha activity with a half-life of ~0.25s. Later results suggest a tentative assignment to 257Lr from the 3n channel with the 246Cm component. No activities assigned to reaction with the 244Cm component have been reported. 249
Bk(18O,αxn)263-xLr (x=3)
This reaction was studied in 1971 by the team at the LBNL in their large study of lawrencium isotopes. They were able to detect an activity assigned to 260Lr. The reaction was further studied in 1988 to study the aqueous chemistry of lawrencium. A total of 23 alpha decays were measured for 260Lr, with a mean energy of 8.03 MeV and an improved half-life of 2.7 minutes. The calculated cross-section was 8.7 nb. 252
Cf(11B,xn)263-xLr (x=5,7??)
This reaction was first studied in 1961 at the University of California by Albert Ghiorso by using a californium target (52% 252Cf). They observed three alpha activities of 8.6 MeV, 8.4 MeV and 8.2 MeV, with half-lives of ~8s and 15s, respectively. The 8.6 MeV activity was tentatively assigned to 257Lr. Later results suggest a reassignment to 258Lr, resulting from the 5n exit channel. The 8.4 MeV activity was also assigned to 257Lr. Later results suggest a reassignment to 256Lr. This is most likely from the 33% 250Cf component in the target rather than from the 7n channel. The 8.2 MeV was subsequently associated with nobelium.
6
Lawrencium
7
252
Cf(10B,xn)262-xLr (x=4,6)
This reaction was first studied in 1961 at the University of California by Albert Ghiorso by using a californium target (52% 252Cf). They observed three alpha activities of 8.6 MeV, 8.4 MeV and 8.2 MeV, with half-lives of ~8s and 15s, respectively. The 8.6 MeV activity was tentatively assigned to 257Lr. Later results suggest a reassignment to 258Lr. The 8.4 MeV activity was also assigned to 257Lr. Later results suggest a reassignment to 256Lr. The 8.2 MeV was subsequently associated with nobelium. 250
Cf(14N,αxn)260-xLr (x=3)
This reaction was studied in 1971 at the LBNL. They were able to identify a 0.7s alpha activity with two alpha lines at 8.87 and 8.82 MeV. This was assigned to 257Lr. 249
Cf(11B,xn)260-xLr (x=4)
This reaction was first studied in 1970 at the LBNL in an attempt to study the aqueous chemistry of lawrencium. They were able to measure a Lr3+ activity. The reaction was repeated in 1976 at Oak Ridge and 26s 256Lr was confirmed by measurement of coincident X-rays. 249
Cf(12C,pxn)260-xLr (x=2)
This reaction was studied in 1971 by the team at the LBNL. They were able to detect an activity assigned to 258Lr from the p2n channel. 249
Cf(15N,αxn)260-xLr (x=2,3)
This reaction was studied in 1971 by the team at the LBNL. They were able to detect an activities assigned to 258Lr and 257Lr from the α2n and α3n and channels. The reaction was repeated in 1976 at Oak Ridge and the synthesis of 258Lr was confirmed. 254
Es + 22Ne - transfer
This reaction was studied in 1987 at the LLNL. They were able to detect new SF activities assigned to 261Lr and 262Lr, resulting from transfer from the 22Ne nuclei to the 254Es target. In addition, a 5 ms SF activity was detected in delayed coincidence with nobelium K X-rays and was assigned to 262No, resulting from the EC of 262Lr.
Synthesis of isotopes as decay products Isotopes of lawrencium have also been identified in the decay of heavier elements. Observations to date are summarised in the table below: Evaporation Residue 267
Bh,
278
263
259
Db
Uut,
274
Lr
Rg,
270
Mt,
266
Bh,
261
Rg,
259
Db
262
Db
258
Lr
257
Db
272
Observed Lr isotope
Lr
268
Mt,
264
Bh,
260
Db
256
Lr
255
Lr
Lawrencium
266
262
261
257
Mt,
Bh,
Bh,
260
8
Bh ,
258
254
Db
Lr
Dbg,m
253
Lrg,m
256
252
Db
Lr
Chronology of isotope discovery Isotope
Year discovered
discovery reaction
252
2001
209
253
1985
209
253
2001
209
254
1985
209
255
1970
243
256
1961? 1965? 1968? 1971
252
257
1958? 1971
249
258
1961? 1971
249
259
1971
248
260
1971
248
261
1987
254
22
Ne
262
1987
254
22
Ne
Lr Lrg Lrm Lr Lr Lr Lr Lr Lr Lr Lr Lr
Bi(50Ti,3n) Bi(50Ti,2n) Bi(50Ti,2n) Bi(50Ti,n) Am(16O,4n) Cf(10B,6n) Cf(15N,α3n) Cf(15N,α2n) Cm(15N,4n) Cm(15N,3n) Es + Es +
Isomerism in lawrencium nuclides 255
Lr
Recent work on the spectroscopy of 255Lr formed in the reaction provided evidence for an isomeric level.
209
Bi(48Ca,2n)255Lr has
253
Lr
A study of the decay properties of 257Db (see dubnium) in 2001 by Hessberger et al. at the GSI provided some data for the decay of 253Lr. Analysis of the data indicated the population of two isomeric levels in 253Lr from the decay of the corresponding isomers in 257Db. The ground state was assigned spin and parity of 7/2-, decaying by emission of an 8794 KeV alpha particle with a half-life of 0.57s. The isomeric level was assigned spin and parity of 1/2-, decaying by emission of an 8722 KeV alpha particle with a half-life of 1.49s.
Lawrencium
9
Chemical yields of isotopes Cold fusion The table below provides cross-sections and excitation energies for cold fusion reactions producing rutherfordium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile 48
Ca
Target 209
Bi
CN
1n
2n
3n
257
Lr
Notes • Los Alamos National Laboratory's Chemistry Division: Periodic Table - Lawrencium
[3]
• Guide to the Elements - Revised Edition, Albert Stwertka, (Oxford University Press; 1998) ISBN 0-19-508083-1
External links • WebElements.com - Lawrencium
[4]
References [1] Nugent, L.J.; Vander Sluis, K.L.; Fricke, Burhard; Mann, J.B.. " Electronic configuration in the ground state of atomic lawrencium (https:/ / kobra. bibliothek. uni-kassel. de/ bitstream/ urn:nbn:de:hebis:34-2008091523764/ 1/ Fricke_electronic_1974. pdf)". Phys. Rev. A 9 (6): 2270–72. . [2] IUPAC "Provisional Recommendations for the Nomenclature of Inorganic Chemistry (2004)" (http:/ / www. iupac. org/ reports/ provisional/ abstract04/ connelly_310804. html) [3] http:/ / periodic. lanl. gov/ elements/ 103. html [4] http:/ / www. webelements. com/ webelements/ elements/ text/ Lr/ index. html
Article Sources and Contributors
Article Sources and Contributors Lawrencium Source: http://en.wikipedia.org/w/index.php?oldid=308809203 Contributors: 03jmgibbens, A new name 2008, Adashiel, Ahoerstemeier, AlimanRuna, ApsbaMd2, Ascidian, Beetstra, Bigtimepeace, Bkell, Bryan Derksen, Burubuz, CJ, Carnildo, Chimneyfreak, Christian List, Conversion script, D6, Dajwilkinson, Darrien, Dave6, David Shay, Dirt palace, Drjezza, Edgar181, Eltomo11, Emperorbma, Enchanter, Enviroboy, Femto, Flying Jazz, Francs2000, Gogo Dodo, Granf, Hak-kâ-ngìn, Hqb, Icairns, Imdwalrus, J.delanoy, J04n, JLaTondre, Jaan513, JackSchmidt, Jamesontai, Jaraalbe, Jiang, John, Karl-Henner, Kgf0, Kingdon, Kingpin13, Kwamikagami, Larryincinci, Lowe4091, Mav, Mdf, Metal Militia, Mimihitam, Naddy, Nergaal, Nihiltres, Nikai, PeterJeremy, Philip Trueman, Physchim62, Polonium, Poolkris, Pras, Psyche825, Remember, Roberta F., Roentgenium111, Romanskolduns, Santăr, Saperaud, Schneelocke, Simon12, Sinnyg, Sl, Snclucas, Squids and Chips, Srnec, Steinbach, Stifynsemons, Stone, Stratocracy, Tagishsimon, Tetracube, Uannis, Vsmith, Werdan7, Yekrats, Zephyrus67, Zoicon5, 142 anonymous edits
Image Sources, Licenses and Contributors image:Lr-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Lr-TableImage.png License: GNU Free Documentation License Contributors: User:Ahoerstemeier
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10
Rutherfordium
1
Rutherfordium lawrencium ← rutherfordium → dubnium Hf ↑ Rf ↓ (Upq) Periodic Table Extended Periodic Table General Name, symbol, number Element category Group, Period, Block Standard atomic weight Electron configuration Electrons per shell
rutherfordium, Rf, 104 transition metals 4, 7, d −1
[267] g·mol 14
[Rn] 5f
2
2, 8, 18, 32, 32, 10, 2 Physical properties
Phase
Unknown Atomic properties
Oxidation states
4 Miscellaneous
CAS registry number
2
6d 7s
53850-36-5 Most-stable isotopes
Rutherfordium
2
Main article: Isotopes of rutherfordium iso 267
Rf
263m
Rf
NA
half-life
DM
syn
1.3 h
SF
syn
~15 m
SF α
263g
syn
8s
SF
262
syn
2.1 s ? 50 ms ?
SF
261m
Rf
syn
1.1 m
261g
Rf
syn
3.0 s
Rf
Rf
α
260
syn
20 ms
SF
259
syn
3.1 s
93% α
Rf
DP
7.90 ?
8.28
257
8.51
257
8.87,8.77
255
No
91% SF 9% α
Rf
DE (MeV)
No
No
7% SF syn
13 ms
SF
257m
Rf
syn
4.0 s
α
9.02,8.97
253
257g
Rf
syn
3.5 s
89% α
8.90,8.78,8.52,8.28
253
258
Rf
11% ε 256
Rf
255
Rf
syn
syn
6.2 ms
1.8 s
No No
257
Lr
99.7% SF 0.3% α
8.79
252
~50% α
8.81,8.77,8.74,8.71
251
No No
~50% SF 254
syn
0.022 ms
SF
253
syn
0.048 ms
SF
Rf Rf
References
Rutherfordium (pronounced /ˌrʌðərˈfɔrdiəm/) is a chemical element in the periodic table that has the symbol Rf and atomic number 104. This is a radioactive synthetic element whose most stable known isotope is 267Rf with a half-life of approximately 1.3 hours. Chemistry experiments have confirmed that rutherfordium behaves as the heavier homologue to hafnium in group 4 (see below).
Rutherfordium
3
History Discovery Element 104 was reportedly first detected in 1966 at the Joint Institute of Nuclear Research at Dubna (then in U.S.S.R.). Researchers there bombarded 242Pu with accelerated 22Ne ions and separated the reaction products by gradient thermochromatography after conversion to chlorides by interaction with ZrCl4. The team identified a spontaneous fission activity contained within a volatile chloride portraying eka-hafnium properties. Although a half-life was not accurately determined, later calculations indicated that the product was most likely 259Rf:[1] 24294Pu
+
2210Ne
→
264-x104Rf
→
264-x104RfCl
4
In 1969 researchers at the University of California, Berkeley conclusively synthesized the element by bombarding a californium-249 target with carbon-12 ions and measured the alpha decay of 257104, correlated with the daughter decay of 253102. [2] 24998Cf
+
126C
→
257104Rf
+4n
The American synthesis was independently confirmed in 1973 and secured the identification of element 104 as the parent by the observation of K X-rays in the elemental signature of the daughter 253No. [3]
Naming controversy The Russian scientists proposed the name Kurchatovium for the new element. The American scientists proposed the name Rutherfordium for the new element. In 1992 the IUPAC/IUPAP Transfermium Working Group (TWG) assessed the claims of discovery and concluded that both teams provided contemporaneous evidence to the synthesis of element 104 and that credit should be shared between the two groups.[1] The American group wrote a scathing response to the findings of the TWG, stating that they had given too much emphasis on the results from the Dubna group. In particular they pointed out that the Russian group had altered the details of their claims several times over a period of 20 years, a fact which the Russian team do not deny. They also stressed that the TWG had given too much credence to the chemistry experiments performed by the Russians and accused the TWG of not having appropriately qualified personnel on the committee. The TWG responded by saying that this was not the case and having assessed each point raised by the American group said that they found no reason to alter their conclusion regarding priority of discovery. [4] It should be said that IUPAC finally used the name suggested by the American team (rutherfordium) which may in some way reflect a change of opinion.[5] As a consequence of the initial competing claims of discovery, an element naming controversy arose. Since the Soviets claimed to have first detected the new element they suggested the name kurchatovium, Ku, in honor of Igor Vasilevich Kurchatov (1903-1960), former head of Soviet nuclear research. This name had been used in books of the Soviet Bloc as the official name of the element. The Americans, however, proposed rutherfordium (Rf) for the new element to honor Ernest Rutherford, who is known as the "father" of nuclear physics. The International Union of Pure and Applied Chemistry (IUPAC) adopted unnilquadium , Unq, as a temporary, systematic element name, derived from the Latin names for digits 1, 0, and 4. In 1994, IUPAC suggested the name dubnium to be used since
Rutherfordium
4
rutherfordium was suggested for element 106 and IUPAC felt that the Dubna team should be rightly recognised for their contributions. However, there was still a dispute over the names of elements 104-107. However in 1997 the teams involved resolved the dispute and adopted the current name rutherfordium.[5]
Newer discoveries Gist of: • isotope synthesis history • nuclear properties • chemical reactions
Future experiments The team at GSI are planning to perform first detailed spectroscopic studies on the isotope 259 Rf. It will be produced in the new reaction: 23892U
+
2412Mg
→
259104Rf
+3n
Isotopes and nuclear properties Nucleosynthesis Chronology of isotope discovery Isotope
Discovered
Reaction [6]
253
1994
204
254
1994
206
255
1974? 1985
207
256
1974? 1985
208
257
1969
249
[2]
258
1969
249
[2]
259
1969
249
[2]
260
1969
248
261
1970
248
261
1996
208
262
1996
244
263
1990?
248
263
2004
248
264
unknown
265
unknown
266
2006?
Rf Rf Rf Rf Rfg,m Rf Rf Rf Rfm Rfg Rf Rfm Rfg Rf Rf Rf
Pb(50Ti,n)
[6]
Pb(50Ti,2n) Pb(50Ti,2n) Pb(50Ti,2n) Cf(12C,4n) Cf(13C,4n) Cf(13C,3n)
Cm(16O,4n) Cm(18O,5n) Pb(70Zn,n)
[7]
[8]
Pu(22Ne,4n)
[9]
Cm(18O,3n) Cm(26Mg,3n)
237
Np(48Ca,3n)
[10]
[11]
Rutherfordium
5
267
2003/2004
238
268
2003?
243
Rf Rf
U(48Ca,3n)
[8]
Am(48Ca,3n)
[12]
Cold fusion 208
Pb(50Ti,xn)258-xRf (x=1,2,3)
This reaction was first studied in 1974 by the team at Dubna. They measured a spontaneous fission activity assigned to 256Rf. [13] The reaction was further studied in 1985 by the GSI team who measured the decay properties of the isotopes 257Rf and 256Rf. The team were able to determine some initial spectroscopic properties of 257Rf and found that the alpha decay pattern was very complicated. [14] After an upgrade of their facilities, they repeated the reaction in 1994 with much higher sensitivity and detected some 1100 atoms of 257Rf and 1900 atoms of 256Rf along with 255Rf in the measurement of the 1n,2n and 3n excitation functions. The large amount of decay data for 257Rf allowed the detection of an isomeric level and the construction of a partial decay level structure which confirmed the very complicated alpha decay pattern. They also found evidence for an isomeric level in 255Rf. [6] The GSI team continued in 2001 with the measurement of the 3n excitation function. In 2002, scientists at the Argonne University in Illinois began their first studies of translawrencium elements with the synthesis and alpha-gamma spectroscopy of 257Rf.[15] In 2004, the GSI began their spectroscopic studies of the 257Rf isotope. 207
Pb(50Ti,xn)257-xRf (x=2)
This reaction was first studied in 1974 by the team at Dubna. They measured a spontaneous fission activity assigned to 255Rf. The reaction was further studied in 1985 by the GSI team who measured the decay properties of the isotope 255Rf. A further spectroscopic study was reported in 2000 which led to a first decay level scheme for the isotope.[16] The isomeric level proposed in 1994 was not found. In 2006, the spectroscopy was continued and the decay scheme was confirmed and improved. [17] 206
Pb(50Ti,xn)256-xRf (x=1,2)
The team at GSI first studied this reaction in 1994 in an effort to study neutron deficient isotopes of rutherfordium. They were able to detect 255Rf and 144 atoms of the new isotope 254Rf, which decayed by spontaneous fission.[6] 204
Currently suggested decay level scheme for 255 Rf from the study reported in 2007 by Hessberger et al. at GSI
Pb(50Ti,xn)254-xRf (x=1)
The team at GSI first studied this reaction in 1994 in an effort to study neutron deficient isotopes of rutherfordium. They were able to detect 14 atoms of the new isotope 253Rf, which decayed by spontaneous fission.[6] 208
Pb(48Ti,xn)256-xRf (x=1)
Rutherfordium In 2006, as part of a program looking at the effect of isospin on the mechanism of cold fusion, the team at LBNL studied this reaction. They measured the 1n excitation function and determined that the change of a Ti-50 projectile to a Ti-48 one significantly reduced the yield, in agreement with predictions. [18] 124
Sn(136Xe,xn)260-xRf
In an important study, in May 2004, the team at GSI attempted the symmetric synthesis of rutherfordium by attempting to fuse two fission fragments. Theory suggests that there may be an enhancement of the yield. No product atoms were detected and a limit of 1000 pb was calculated. Hot fusion 238
U(26Mg,xn)264-xRf (x=3,4,5,6)
The hot fusion reaction using a uranium target was first reported in 2000 by Yuri Lazarev and the team at the Flerov Laboratory of Nuclear Reactions (FLNR). They were able to observe decays from 260Rf and 259Rf in the 4n and 5n channels. [19] They measured yields of 240 pb in the 4n channel and 1.1 nb in the 5n channel. In 2006, as part of their program on the study of uranium targets in hot fusion reactions, the team at LBNL measured the 4n,5n and 6n excitation functions for this reaction and observed 261Rf in the 3n exit channel.[20] [21] 244
Pu(22Ne,xn)266-xRf (x=4,5)
This reaction was reported in 1996 at LBNL in an attempt to study the fission characteristics of 262Rf. The team were able to detect the SF of 262Rf and determine its half-life as 2.1 s, in contrast to earlier reports of a 47 ms activity. It was suggested that the two half-lives may be related to different isomeric states. [9] The reaction was further studied in 2000 by Yuri Lazarev and the team at Dubna. They were able to observe 69 alpha decays from 261Rf and spontaneous fission of 262Rf. [22] Later work on hassium has allowed a reassignment of the 5n product to 261mRf. 242
Pu(22Ne,xn)264-xRf (x=3,4?,5?)
The synthesis of element 104 was first attempted in 1964 by the team at Dubna using this reaction. The first study produced evidence for a 0.3 s spontaneous fission (SF) activity tentatively assigned to 260104 or 259104 and an unidentified 8 s SF activity. The former activity was later retracted and the latter activity associated with the now-known 259104.[1] In 1966, in their discovery experiment, the team repeated the reaction using a chemical study of volatile chloride products. The group was able to identify a volatile chloride decaying by short spontaneous fission with eka-hafnium properties. This gave strong evidence for the formation of [104]Cl4 and the team suggested the name kurchatovium. Although a half-life was not accurately measured, later evidence suggested that the product was most likely 259104. [1] In 1968, the team searched for alpha decay from 260104 but were unable to detect such activity. In 1970, the team repeated the reaction once again and confirmed the ~0.2 s SF activity. They also repeated the chemistry experiment and obtained identical results to their 1966 experiment and calculated a likely half-life of ~0.5 seconds for the SF activity. In 1971, the reaction was repeated again and 0.1 s and 4.5 s SF activities were found. The 4.5 s activity was correctly assigned to 259104.[1] A chemistry experiment in the same year reaffirmed the formation of a 0.3 SF activity for an eka-hafnium product.[1] Later, the 0.1-0.3 s SF activity was retracted as belonging to a kurchatovium isotope but the observation of eka-hafnium reactivity remained and was the
6
Rutherfordium basis of their successful claim to discovery.[1] The reaction was further studied in 2000 by Yuri Lazarev at Dubna. They were able to observe 261Rf in the 3n channel, later reassigned to 261mRf. 242
Pu(20Ne,xn)262-xRf
This reaction was first studied in 1964 to assist in the assignments using the analogous reaction with a Ne-22 beam. The Dubna team were unable to detect any 0.3 s spontaneous fission activities.[1] The reaction was later studied in 2003 at the Paul Scherrer Institute (PSI) in Bern, Switzerland. They detected some spontaneous fission activities but were unable to confirm the formation of 259Rf.[23] 248
Cm(22Ne,αxn)266-xRf (x=3?)
This reaction was studied in 1999 at the University of Bern, Switzerland in order to search for the new isotope 263Rf. A rutherfordium fraction was separated and several SF events with long lifetimes and alpha decays with energy 7.8 MeV and 7.9 MeV were observed. A second experiment using a study of the fluoride of rutherfordium products also produced 7.9 MeV alpha decays.[24] 248
Cm(18O,xn)266-xRf (x=3?,5)
This reaction was first studied in 1970 by Albert Ghiorso at the Lawrence Berkeley National Laboratory (LBNL). The team identified 261Rf in the 5n channel using the method of correlation of genetic parent-daughter decays. A half-life of 65 s was determined. [7] A repeat later that year using cation exchange chromatography indicated that the product did not form a +2 or +3 cation and behaved as eka-hafnium. A study of the properties of rutherfordium isotopes was performed in 1981 at the LBNL. In a series of reactions, a 1.5 s SF activity was identified and assigned to a fermium descendant although later evidence indicates a possible assignment to 262Rf. In contrast, in a subsequent review of isotope properties by Somerville et al. at LBNL in 1985, a 47 ms SF activity was assigned to 262Rf. This assignment has not been verified. [25] The reaction was further studied in 1991 by Czerwinski et al. at the LBNL. In this experiment, spontaneous fission activities with long lifetimes were observed in rutherfordium fractions and tentatively assigned to 263Rf. In 1996, chemical studies on the chloride of rutherfordium was published by the LBNL. In this experiment, the half-life was improved to 78 s. A repeat of the experiment in 2000 assessing the volatility of the bromide further refined the half-life to 75 s. 248
Cm(16O,xn)264-xRf (x=4)
This reaction was studied in 1969 by Albert Ghiorso at the University of California. The aim was to detect the 0.1-0.3 s SF activity reported at Dubna, assigned to 260104. They were unable to do so, only observing a 10-30 ms SF activity, correctly assigned to 260104. The failure to observe the 0.3 s SF activity identified by Dubna gave the Americans the incentive to name this element rutherfordium.[1] 246
Cm(18O,xn)264-xRf
In an attempt to unravel the properties of spontaneous fission activities in the formation of rutherfordium isotopes, this reaction was performed in 1976 by the FLNR. They observed an 80 ms SF activity. Subsequent work led to the complete retraction of the 0.3s - 0.1s - 80 ms SF activities observed by the Dubna team and associated with background signals.[1] 249
Bk(15N,xn)264-xRf (x=4)
This reaction was studied in 1977 by the team in Dubna. They were able to confirm the detection of a 76 ms SF activity. The assignment to rutherfordium isotopes was later
7
Rutherfordium
8
retracted. The LBNL re-studied the reaction in 1980 and in 1981 they reported that they were unable to confirm the ~80 ms SF activity. The Dubna team were able to measure a 28 ms SF activity in 1985 and assigned the isotope correctly to 260104.[1] 249
Cf(13C,xn)262-xRf (x=4)
This use of californium-249 as a target was first studied by Albert Ghiorso and the team at the University of California in 1969. They were able to observe a 11 ms SF activity which they correctly assigned to 258104.[2] 249
Cf(12C,xn)261-xRf (x=3,4)
In their 1969 discovery experiments, the team at University of California also used a C-12 beam to irradiate a Cf-249 target. They were able to confirm the 11 ms SF activity found with a C-13 beam and again correctly assigned to 258104. The actual discovery experiment was the observation of alpha decays genetically linked to 253102 and therefore positively identified as 257104.[2] In 1973, Bemis and his team at Oak Ridge confirmed the discovery by measuring coincident X-rays from the daughter 253 102.[3] As decay product Isotopes of rutherfordium have also been identified
Currently suggested decay level scheme for 257 g,m Rf from the study performed in 2004 [6] by Hessberger et al. at GSI
in the decay of heavier elements. Observations to date are summarised in the table below: Evaporation Residue 288
Observed Rf isotope 268
115
Rf (possible EC of
291
116 ,
287
114 ,
283
267
112
Rf
282
266
271
263g
263
263m
113
Rf (EC of
Hs
Rf (EC of
266
Sg (possibly 112 ,
273
Ds ,
266m
262
Sg)
269
Hs ,
265
Sg
261m
Rf ,
267
263
259
269
265
261
257
264
260
Ds , Hs ,
259
Sg
Db)
Hs , Hs , Sg
Sg Sg
263
Rf (possibly
271
Ds ,
266
Rf
Db
277
268
Rf Rf
256
Rf
255
Rf
261
Rf
Db)
262m
Rf)
Db)
Rutherfordium Unconfirmed and retracted isotopes 268
Rf
In the synthesis of ununpentium, the isotope 288115 has been observed to decay to 268Db which undergoes spontaneous fission with a half life of 29 hours. Given that the electron capture of 268Db cannot be detected, these SF events may in fact be due to the SF of 268Rf, in which case the half-life of this isotope cannot be extracted. [12] 266
Rf
In the synthesis of ununtrium, the isotope 282113 has been observed to decay to 266Db which undergoes spontaneous fission with a half life of 22 minutes. Given that the electron capture of 266Db cannot be detected, these SF events may in fact be due to the SF of 266Rf, in which case the half-life of this isotope cannot be extracted. [11] 265
Rf
In 1999, American scientists at the University of California, Berkeley, announced that they has succeeded in synthesizing three atoms of 293118. These parent nuclei successively emitted seven alpha particles to form 265Rf nuclei. Their claim was retracted in 2001. As such, this rutherfordium isotope is unconfirmed or unknown.[26] 255m
Rf
A detailed spectroscopic study of the production of 255Rf nuclei using the reaction 206 Pb(50Ti,n)255Rf allowed the tentative identification of an isomeric level in 255Rf. A more detailed study later confirmed that this was not the case.
Nuclear isomerism 263
Rf
Initial work on the synthesis of rutherfordium isotopes by hot fusion pathways focused on the synthesis of 263Rf. Several studies have indicated that this nuclide decays primarily by spontaneous fission with a long half-life of 10–20 minutes. Alpha particles with energy 7.8-7.9 MeV have also been associated with this nucleus. More recently, a study of hassium isotopes allowed the synthesis of an atom of 263Rf decaying by spontaneous fission with a short half-life of 8 seconds. These two different decay modes must be associated with two isomeric states. Specific assignments are difficult due to the low number of observed events. It is reasonable to tentatively assign the long life to a meta-stable state, namely 263m Rf, and the shorter life to the ground state, namely 263gRf. Further studies are required to allow a definite assignment. 261
Rf
Early research on the synthesis of rutherfordium isotopes utilised the 244Pu(22Ne,5n)261Rf reaction. The product was found to undergo exclusive 8.28 MeV alpha decay with a half-life of 78 seconds. Later studies by the GSI team on the synthesis of element 112 and hassium isotopes produced conflicting data. In this case, 261Rf was found to undergo 8.52 MeV alpha decay with a short half-life of 4 seconds. Later results indicated a predominant fission branch. These contradictions led to some doubt on the discovery of element 112. However, it is now understood that the first nucleus belongs to an isomeric meta-stable state, namely 261m Rf and the latter to the ground state isomer, namely 261gRf. [27] The discovery and confirmation of 261gRf provided proof for the discovery of ununbium in 1996. 257
Rf
9
Rutherfordium
10
A detailed spectroscopic study of the production of 257Rf nuclei using the reaction 208 Pb(50Ti,n)257Rf allowed the identification of an isomeric level in 257Rf. The work confirmed that 257gRf has a very complicated spectrum with as many as 15 alpha lines. A level structure diagram was calculated for both isomers.
Chemical yields of isotopes Cold fusion The table below provides cross-sections and excitation energies for cold fusion reactions producing rutherfordium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
CN
1n
2n
50
208
258
38.0 nb , 17.0 MeV
50
207
257
4.8 nb
50
206
256
800 pb , 21.5 MeV
50
204
254
190 pb , 15.6 MeV
48
208
256
380 pb , 17.0 MeV
Ti
Pb
Ti
Rf
Pb
Ti
Rf
Pb
Ti
Rf
Pb
Ti
Rf
Pb
Rf
12.3 nb , 21.5 MeV
3n 660 pb , 29.0 MeV
2.4 nb , 21.5 MeV
Hot fusion The table below provides cross-sections and excitation energies for hot fusion reactions producing rutherfordium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
CN
3n
4n
5n
26
238
264
240 pb
1.1 nb
22
244
266
+
4.0 nb
18
248
266
+
13.0 nb
Mg Ne O
U Pu Cm
Rf Rf Rf
Chemical properties Electronic structure Rutherfordium is element 104 in the Periodic Table. The two forms of the projected electronic structure are: Bohr model Quantum mechanical model
2, 8, 18, 32, 32, 10, 2 1s2 2s22p6 3s23p6 4s23d104p6 5s24d105p6 6s24f145d106p6 7s25f146d2
Rutherfordium
11
Extrapolated properties Oxidation states Element 104 is projected to be the first member of the 6d series of transition metals and the heaviest member of group IV in the Periodic Table, below titanium, zirconium and hafnium. It may also be named eka-hafnium or dvi-zirconium and some of its properties may be extrapolated from the group trends of elements. The IV oxidation state is the only stable state for the latter two elements and therefore rutherfordium should also portray a stable +4 state. Chemistry In an analogous manner to zirconium and hafnium, rutherfordium is projected to form a very stable, high melting point oxide, RfO2. It should also react with halogens to form tetrahalides, RfX4, which hydrolyse on contact with water to form oxyhalides RfOX2. The tetrahalides should be volatile solids existing as monomeric tetrahedral molecules in the vapour phase. In the aqueous phase, the Rf4+ ion should hydrolyse less than titanium(IV) and to a similar extent to zirconium and hafnium, thus leading to the rutherfordyl oxyion, RfO22+. Treatment of the halides with halide ions promotes the formation of complex ions. The use of chloride and bromide ion should form the hexahalide complexes RfCl62− and RfBr62−. For the fluoride complexes, zirconium and hafnium tend to form hepta- and octacomplexes. Thus, for the larger rutherfordium ion, the complexes RfF62−, RfF73− and RfF84− are possible.
Experimental chemistry Summary of compounds and complex ions Formula
Names(s)
RfCl4
rutherfordium tetrachloride ; rutherfordium(IV) chloride
RfBr4
rutherfordium tetrabromide ; rutherfordium(IV) bromide
RfOCl2
rutherfordium oxychloride ; rutherfordyl(IV) chloride; rutherfordium(IV) dichloride oxide
[RfCl6]2-
hexachlororutherfordate(IV)
[RfF6]2−
hexafluororutherfordate(IV)
K2[RfCl6]
potassium hexachlororutherfordate(IV)
Gas phase Early work on the study of the chemistry of rutherfordium focused on gas thermochromatography and measurement of relative deposition temperature adsorption curves. The initial work was carried out at Dubna in an attempt to reaffirm their discovery of the element. Recent work is more reliable regarding the identification of the parent rutherfordium radioisotopes. The isotope 261mRf has been used for these studies. The experiments relied on the expectation that rutherfordium would begin the new 6d series of elements and should therefore from a volatile tetrachloride due to the tetrahedral nature of the molecule:
Rutherfordium
12 As series of experiments have confirmed that rutherfordium behaves as a typical member of group 4 forming a tetravalent chloride (RfCl4) and bromide (RfBr4) as well as an oxychloride (RfOCl2).[28] Aqueous phase
Rutherfordium is expected to have the electron configuration [Rn]5f14 6d2 7s2 and therefore behave as the heavier homologue of hafnium in group 4 of the Periodic Table. It should therefore readily form a 4+ hydrated Rf ion in strong acid solution and should readily form complexes in hydrochloric acid, hydrobromic or hydrofluoric acid solutions. The tetrahedral molecule RfCl4
The most conclusive aqueous chemistry studies of rutherfordium have been performed by the Japanese team at JAERI using the radioisotope 261mRf. Extraction experiments from hydrochloric acid solutions using isotopes of rutherfordium, hafnium, zirconium and thorium have proved a non-actinide behaviour. A comparison with its lighter homologues placed rutherfordium firmly in group 4 and indicated the formation of a hexachlororutherfordate complex in chloride solutions, in a manner similar to hafnium and zirconium. [29] 261m
Rf4+ + 6 Cl− → [261mRfCl6]2−
Very similar results were observed in hydrofluoric acid solutions. Differences in the extraction curves were interpreted as a weaker affinity for fluoride ion and the formation of the hexafluororutherfordate ion, whereas hafnium and zirconium ions complex seven or eight fluoride ions at the concentrations used: 261m
Rf4+ + 6 F− → [261mRfF6]2−
References [1] Barber, R. C.; Greenwood, N. N.; Hrynkiewicz, A. Z.; Jeannin, Y. P.; Lefort, M.; Sakai, M.; Ulehla, I.; Wapstra, A. P.; Wilkinson, D. H. (1993). " Discovery of the transfermium elements. Part II: Introduction to discovery profiles. Part III: Discovery profiles of the transfermium elements (http:/ / www. iupac. org/ publications/ pac/ 65/ 8/ 1757/ )". Pure and Applied Chemistry 65 (8): 1757–1814. doi: 10.1351/pac199365081757 (http:/ / dx. doi. org/ 10. 1351/ pac199365081757). . [2] Ghiorso, A.; Nurmia, M.; Harris, J.; Eskola, K.; Eskola, P. (1969). " Positive Identification of Two Alpha-Particle-Emitting Isotopes of Element 104 (http:/ / prola. aps. org/ abstract/ PRL/ v22/ i24/ p1317_1)". Physical Review Letters 22: 1317–1320. doi: 10.1103/PhysRevLett.22.1317 (http:/ / dx. doi. org/ 10. 1103/ PhysRevLett. 22. 1317). . [3] Bemis, C. E. C. E. Bemis, Jr., Silva, R. J.; Hensley, D. C.; Keller Jr., O. L.; Tarrant, J. R.; Hunt, L. D.; Dittner,P. F.; Hahn, R. L.; Goodman, C. D. (1973). "X-Ray Identification of Element 104". Physical Review Letters 31: 647–650. doi: 10.1103/PhysRevLett.31.647 (http:/ / dx. doi. org/ 10. 1103/ PhysRevLett. 31. 647). [4] Ghiorso, A.; Seaborg, G. T.; Organessian, Yu. Ts.; Zvara, I.; Armbruster, P.; Hessberger, F. P.; Hofmann, S.; Leino, M.; Munzenberg, G.; Reisdorf W.; Schmidt, K.-H. (1993). " Responses on 'Discovery of the transfermium elements' by Lawrence Berkeley Laboratory, California; Joint Institute for Nuclear Research, Dubna; and Gesellschaft fur Schwerionenforschung, Darmstadt followed by reply to responses by the Transfermium Working Group (http:/ / www. iupac. org/ publications/ pac/ 1993/ pdf/ 6508x1815. pdf)". Pure and Applied Chemistry 65 (8): 1815–1824. doi: 10.1351/pac199365081815 (http:/ / dx. doi. org/ 10. 1351/ pac199365081815). . [5] " Names and symbols of transfermium elements (IUPAC Recommendations 1997) (http:/ / www. iupac. org/ publications/ pac/ 1997/ pdf/ 6912x2471. pdf)". Pure and Applied Chemistry 69 (12): 2471–2474. 1997. doi: 10.1351/pac199769122471 (http:/ / dx. doi. org/ 10. 1351/ pac199769122471). . [6] Heßberger, F. P.; Hofmann, S.; Ninov, V.; Armbruster, P.; Folger, H.; Münzenberg, G.; Schött, H. J.; Popeko, A. K.; Yeremin, A. V. ; Andreyev, A. N.; Saro, S. (1997). "Spontaneous fission and alpha-decay properties of neutron deficient isotopes 257-253104 and 258106". Zeitschrift für Physik a Hadrons and Nuclei 359: 415. doi: 10.1007/s002180050422 (http:/ / dx. doi. org/ 10. 1007/ s002180050422).
Rutherfordium [7] Ghiorso, A.; Nurmia, M.; Eskola, K.; Eskola P. (1970). "261Rf; new isotope of element 104". Physics Letters B 32 (2): 95–98. doi: 10.1016/0370-2693(70)90595-2 (http:/ / dx. doi. org/ 10. 1016/ 0370-2693(70)90595-2). [8] see ununbium [9] M. R. Lane, K. E. Gregorich, D. M. Lee, M. F. Mohar, M. Hsu, C. D. Kacher, B. Kadkhodayan, M. P. Neu, N. J. Stoyer, E. R. Sylwester, J. C. Yang, and D. C. Hoffman (1996). "Spontaneous fission properties of 104262Rf". Physical Review C 53 (6): 2893–2899. doi: 10.1103/PhysRevC.53.2893 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 53. 2893). [10] see hassium [11] see ununtrium [12] see ununpentium [13] Oganessian, Yu.Ts.; Demin, A. G.; Il'inov, A. S.; Tret'yakova, S. P.; Pleve, A. A.; Penionzhkevich, Yu. É.; Ivanov M. P.; Tret'yakov, Yu. P. (1975). " Experiments on the synthesis of neutron-deficient kurchatovium isotopes in reactions induced by 50Ti Ions (http:/ / www. springerlink. com/ content/ m87v320736564810/ )". Nuclear Physics A 38 (6): 492–501. doi: 10.1016/0375-9474(75)91140-9 (http:/ / dx. doi. org/ 10. 1016/ 0375-9474(75)91140-9). . [14] Heßberger, F. P.; Münzenberg, G.; Hofmann, S.; Reisdorf, W.; Schmidt, K. H.; Schött, H. J.; Armbruster, P.; Hingmann, R.; Thuma, B.; Vermeulen, D. (1985). "Study of evaporation residues produced in reactions of 207 208 , Pb with 50Ti". Zeitschrift für Physik a Atoms and Nuclei 321: 19–26. doi: 10.1007/BF01493453 (http:/ / dx. doi. org/ 10. 1007/ BF01493453). [15] Qian, J.; Heinz, A.; Winkler, R.; Vinson, J.; Janssens, R. V. F.; Peterson, D.; Seweryniak, D.; Back, B.; Carpenter, M. P.; Savard, G.; Hecht, A. A.; Jiang, C. L.; Khoo, T. L.; Kondev, F. G.; Lauritsen, T.; Lister, C. J.; Robinson, A.; Wang, X.; Zhu, S.; Gansworthy, A. B.; Asai, M.. " Alpha decay of 257Rf (http:/ / adsabs. harvard. edu/ abs/ 2007APS. . DNP. JG010Q)". American Physical Society, 2007 Annual Meeting of the Division of Nuclear Physics: JG.010. . [16] Heßberger, F.P.; Hofmann, S.; Ackermann, D.; Ninov, V.; Leino, M.; Münzenberg, G.; Saro, S.; Lavrentev, A.; Popeko, A.G.; Yeremin, A.V.; Stodel, Ch. (2001). "[http://www.edpsciences.org/articles/epja/abs/2001/09/epja1103/epja1103.html Decay properties of neutron-deficient isotopes 256,257Db, 255Rf, 252,253Lr"]]". European Physical Journal A 12: 57–67. http:/ / www. edpsciences. org/ articles/ epja/ abs/ 2001/ 09/ epja1103/ epja1103. html. [17] Heßberger, F.P.; Hofmann, S.; Ackermann, D.; Antalic, S.; Kindler, B.; Kojouharov, I.; Kuusiniemi, P.; Leino, M.; Lommel, B.;Mann, R.; Nishio K.; Popeko, A. G.; Sulignano B.; Saro, S.; Streicher, B.; Venhart, M.; Yeremin, A. V. (2006). " Alpha-gamma decay studies of 255Rf, 251No and 247Fm (http:/ / www. edpsciences. org/ articles/ epja/ abs/ 2006/ 15/ 10050_2006_Article_100239/ 10050_2006_Article_100239. html)". European Physical Journal A 30: 561–569. . [18] Dragojević, I.; Gregorich, K. E.; Düllmann, Ch.E.; Garcia, M.A. (2005). " Measurement of 208Pb(48Ti,n)255Rf excitation function (http:/ / rnc. lbl. gov/ nsd/ annualreport2005/ contributions/ Dragojevic_LE. pdf)". LBNL annual report. . Retrieved 2008-02-29. [19] Lazarev, Yu. A.; Lobanov, Yu. V.; Oganessian, Yu. Ts.; Utyonkov, V. K.; Abdullin, F. Sh.; Polyakov, A. N.; Rigol, J.; Shirokovsky, I. V.; Tsyganov, Yu. S.; Iliev, S.; Subbotin, V. G.; Sukhov, A. M.; Buklanov, G. V.; Mezentsev, A. N.;Subotic, K.; Moody, K. J.; Stoyer, N. J.; Wild, J. F.; Lougheed, R. W. (2000). " Decay properties of 257No, 261Rf, and 262Rf (http:/ / prola. aps. org/ abstract/ PRC/ v62/ i6/ e064307)". Physical Review C 62 (6): 064307. doi: 10.1103/PhysRevC.62.064307 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 62. 064307). . [20] Gregorich, K.E.; Ch.E. Düllmann, C.M. Folden III, R. Sudowe, S.L. Nelson, J.M. Gates, I. Dragojević, M.A. Garcia, Y.H. Chung, R. Eichler, G.K. Pang, A. Türler, A. Yakushev, D.C. Hoffman, H. Nitsche (2005). " Systematic Study of Heavy Element Production in Compound Nucleus Reactions with 238U Targets (http:/ / rnc. lbl. gov/ nsd/ annualreport2005/ contributions/ Gregorich_LE. pdf)". LBNL annual report. . Retrieved 2008-02-29. [21] J. M. Gates, M. A. Garcia, K. E. Gregorich, Ch. E. Düllmann, I. Dragojević, J. Dvorak, R. Eichler, C. M. Folden III, W. Loveland, S. L. Nelson, G. K. Pang, L. Stavsetra, R. Sudowe, A. Türler, and H. Nitsche (2008). "Synthesis of rutherfordium isotopes in the 238U(26Mg,xn)264-xRf reaction and study of their decay properties". Physical Review C 77: 034603. doi: 10.1103/PhysRevC.77.034603 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 77. 034603). [22] Yu. A. Lazarev , Yu. V. Lobanov, Yu. Ts. Oganessian, V. K. Utyonkov, F. Sh. Abdullin, A. N. Polyakov, J. Rigol, I. V. Shirokovsky, Yu. S. Tsyganov, S. Iliev, V. G. Subbotin, A. M. Sukhov, G. V. Buklanov, A. N. Mezentsev, and K. Subotic K. J. Moody, N. J. Stoyer, J. F. Wild, and R. W. Lougheed (1996). "Decay properties of 257No, 261Rf, and 262Rf". Physical Review C 62 (6): 064307. doi: 10.1103/PhysRevC.62.064307 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 62. 064307). [23] " 20Ne On 244Pu - First Preliminary Results (http:/ / lch. web. psi. ch/ pdf/ anrep03/ 03. pdf)". . [24] "An EC-branch in the decay of 27-s 263Db: Evidence for the new isotope 263Rf" (http:/ / www. ulrich-rieth. de/ publikationen/ RCA0301_059. PDF), Kratz et al., GSI Annual report 2001. Retrieved on 2008-02-29
13
Rutherfordium [25] Somerville, L. P.; M. J. Nurmia, J. M. Nitschke, and A. Ghiorso E. K. Hulet and R. W. Lougheed (1985). " Spontaneous fission of rutherfordium isotopes (http:/ / prola. aps. org/ abstract/ PRC/ v31/ i5/ p1801_1)". Physical Review C 31 (5): 1801–1815. doi: 10.1103/PhysRevC.31.1801 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 31. 1801). . [26] see ununoctium 261 [27] "EVIDENCE FOR ISOMERIC STATES IN Rf" (http:/ / lch. web. psi. ch/ pdf/ anrep01/ B-02heavies. pdf), Dressler et al., PSI Annual Report 2001. Retrieved on 2008-01-29 [28] http:/ / lch. web. psi. ch/ pdf/ TexasA& M/ TexasA& M. pdf (Gas Chem 2007 Review) [29] Yuichi Nagame, ; K. Tsukada, ; M. Asai, ; A. Toyoshima, ; K. Akiyama, ; Y. Ishii, ; T. Kaneko- Sato, ; M. Hirata, ; I. Nishinaka, ; S. Ichikawa, ; H. Haba, ; Shuichi Enomoto (2005). " (Rf aqueous chem JAERI) Chemical studies on rutherfordium (Rf) at JAERI (http:/ / wwwsoc. nii. ac. jp/ jnrs/ paper/ JN62/ jn6202. pdf)". Radiochimica Acta 93: 519. doi: 10.1524/ract.2005.93.9-10.519 (http:/ / dx. doi. org/ 10. 1524/ ract. 2005. 93. 9-10. 519). (Rf aqueous chem JAERI).
External links • WebElements.com - Rutherfordium (http:/ / www. webelements. com/ webelements/ elements/ text/ Rf/ index. html)
14
Article Sources and Contributors
Article Sources and Contributors Rutherfordium Source: http://en.wikipedia.org/w/index.php?oldid=300776541 Contributors: Achaemenes, Ahoerstemeier, AlimanRuna, Altenmann, Anclation, Andres, Andy120, Antonio Lopez, Astavats, Bhound89, Bored461, Bryan Derksen, Bubba hotep, Cameron Nedland, Carnildo, Chris the speller, CieloEstrellado, Control.valve, Conversion script, DIG, Darrien, DavidRF, Doug Funny, Dougdp, Drjezza, ESkog, Ecopetition, Edgar181, Emperorbma, Eras-mus, Femto, Gamma, Gbr3, Graham87, Greatpatton, Greenguy746, Gurch, HTait, Hak-kâ-ngìn, Heliomance, Icairns, Ideyal, John, Kalamkaar, Karelj, Karl-Henner, Kelovy, Kgf0, Kingdon, Kiwi137, Kwamikagami, Kwi, Leokor, Mav, Mdf, Mgiganteus1, Miss Madeline, Mortdefides, Nergaal, Nomad, Nrjs, Ostrich11, PlatinumX, Poolkris, Pras, Proteus, Psyche825, Qwerty Binary, Rednblu, Remember, RexNL, Rifleman 82, Roentgenium111, Romanm, Romanskolduns, Ryoung122, Saperaud, Sceptre, Schneelocke, Seth Ilys, Sfisher, Sharkface217, Sl, Soliloquial, Stifynsemons, Stone, Sverre, Tagishsimon, Tetracube, Walkerma, Warut, Yekrats, Yms, 92 anonymous edits
Image Sources, Licenses and Contributors image:Rf-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Rf-TableImage.png License: GNU Free Documentation License Contributors: Daniel Mayer (w:en:Maveric149) and/or Arnaud Gaillard (Utilisateur:GreatpattonGreatpatton) Image:255Rf decay scheme 2006.png Source: http://en.wikipedia.org/w/index.php?title=File:255Rf_decay_scheme_2006.png License: Public Domain Contributors: User:Drjezza Image:257Rf decay scheme 2006.png Source: http://en.wikipedia.org/w/index.php?title=File:257Rf_decay_scheme_2006.png License: Public Domain Contributors: User:Drjezza Image:RfCl4.png Source: http://en.wikipedia.org/w/index.php?title=File:RfCl4.png License: Public Domain Contributors: User:Drjezza
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
15
Dubnium
1
Dubnium rutherfordium ← dubnium → seaborgium Ta ↑ Db ↓ (Upp) Periodic Table Extended Periodic Table General Name, symbol, number Element category Group, Period, Block Appearance
dubnium, Db, 105 transition metals 5, 7, d unknown, probably silvery white or metallic gray
Standard atomic weight Electron configuration Electrons per shell
−1
[268] g·mol 14
[Rn] 5f
3
2
6d 7s
2, 8, 18, 32, 32, 11, 2 Physical properties
Phase
presumably a solid Density (near r.t.)
−3
unknown g·cm Atomic properties
Crystal structure Oxidation states
unknown 5
Atomic radius (calc.)
unknown pm
Covalent radius
unknown pm Miscellaneous
CAS registry number
53850-35-4 Most-stable isotopes
Dubnium
2
Main article: Isotopes of dubnium iso 268
Db
NA syn
half-life [1]
16h
DM
DE (MeV)
SF ε?
268
Rf
267
syn
1.2 h
SF
266
syn
22 m
ε
263
syn
27 s
56% SF
Db Db Db
266
Rf
41% α
8.36
3% ε 262
Db
syn
34 s
DP
67% α
259
Lr
263m
Rf
8.66,8.45
258
Lr
33% SF 261
syn
1.8 s
α
8.93
257
260
syn
1.5 s
α
9.13,9.08,9.05
256
259
syn
0.5 s
α
9.47
255
258
syn
4.4 s
67% α
9.17,9.08,9.01
254
Db Db Db Db
33% ε
Lr Lr Lr Lr
258
Rf
257m
Db
syn
0.76 s
α
9.16
253
257g
Db
syn
1.50 s
α
9.07,8.97
253
256
syn
1.6 s
70% α
9.12,9.08,9.01,8.89
253
Db
30% ε
Lr Lr Lr
256
Rf
References
Dubnium (pronounced /ˈduːbniəm/) is a chemical element in the periodic table that has the symbol Db and atomic number 105. This is a radioactive synthetic element whose most stable isotope is 268Db with a half life of 28 hours[1] . An element of with a half-life as long as this would be much easier to work with, than some other elements such as Element 118 which has a half-life of 3 milliseconds. This is the longest lived transactinide isotope and is a reflection of the stability of the Z = 108 and N = 162 closed shells and the effect of odd particles in nuclear decay. Chemistry experiments have provided sufficient evidence to confidently place dubnium in group 5 of the Periodic Table.
Dubnium
3
Discovery profile Element 105 was first reported by Russian scientists, in 1968-1970 at the Joint Institute for Nuclear Research in Dubna, Russia. The 1968 work was based on the detection of correlated decays of element 105 to known daughter nuclei using the reaction 243 Am(22Ne,xn). They reported a 9.40 MeV and a 9.70 MeV alpha-activity and assigned the decays to the isotopes 260105 or 261105. In 1970 they expanded their work by the application of thermal gradient chromatography and detection by spontaneous fission. They observed a 2.2 s SF activity in a fraction portraying niobium-like characteristics and assigned the activity to 261DbCl5. In late April 1970 researchers led by Albert Ghiorso working at the University of California, Berkeley published a convincing synthesis of 260Db in the reaction: 24998Cf
+
157N
→
260105Db
+4n
The team claimed that 260Db decayed by 9.10 MeV alpha-emission with a half-life of 1.6 seconds to 256Lr. Decay data for 256Lr agreed with the literature values and provided strong support to their claim. These results by the Berkeley scientists did not confirm the Soviet findings regarding the 9.40 MeV or 9.70 MeV alpha-decay of 260Db. In 1971, the Russian team repeated their reaction using an improved set-up and were able to confirm the decay data for 260Db using the reaction: 243
Am +
22
Ne →
260
Db + 5 n
In 1976, the Russian team continued their study of the reaction using thermal gradient chromatography and were able to identify the product 260DbBr5. In 1977, all doubt was dispelled by the L X-ray elemental detection of lawrencium isotopes from the reaction: 249
Cf +
15
N→
260
Db + 4 n
In 1992 the TWG assessed the claims of the two groups and concluded that confidence in the discovery grew from results from both laboratories and the claim of discovery should be shared.[2]
Proposed names Historically element 105 has been called eka-tantalum, reflecting Mendeleev's placeholder terminology. The American team proposed that the new element should be named hahnium (Ha), in honor of the late German physicist Otto Hahn. Consequently this was the name that most American and Western European scientists used and appears in many papers published at the time. The Russian team proposed the name nielsbohrium (Ns) in honor of the Danish nuclear physicist Niels Bohr. An element naming controversy erupted between the two groups. The International Union of Pure and Applied Chemistry (IUPAC) thus adopted unnilpentium (Unp) as a temporary, systematic element name. Attempting to resolve the issue, in 1994, the IUPAC proposed the name joliotium (Jl), after the French physicist Frédéric Joliot-Curie. The two principal claimants still disagreed about the names of elements 104-106. However in 1997 they
Dubnium
4
resolved the dispute and adopted the current name, dubnium (Db), after the city that contains the Russian Joint Institute for Nuclear Research. It was argued by IUPAC that the Berkeley laboratory had already been recognized several times in the naming of elements (i.e., berkelium, californium, americium) and that the acceptance of the names rutherfordium and seaborgium for elements 104 and 106 should be offset by recognizing the contributions of the Russian team to the discovery of elements 104,105 and 106.[3] [4]
Electronic structure Dubnium is element 105 in the Periodic Table. The two forms of the projected electronic structure are: Bohr model Quantum mechanical model
2, 8, 18, 32, 32, 11, 2 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f146d3
Extrapolated chemical properties of eka-tantalum/dvi-niobium Oxidation states Element 105 is projected to be the second member of the 6d series of transition metals and the heaviest member of group V in the Periodic Table, below vanadium, niobium and tantalum. All the members of the group readily portray their oxidation state of +5 and the state becomes more stable as the group is descended. Thus dubnium is expected to form a stable +5 state. For this group, +4 and +3 states are also known for the heavier members and dubnium may also form these reducing oxidation states.
Chemistry In an extrapolation of the chemistries from niobium and tantalum, dubnium should react with oxygen to form an inert pentoxide, Db2O5. In alkali, the formation of an orthodubnate complex, DbO43−, is expected. Reaction with the halogens should readily form the pentahalides, DbX5. The pentachlorides of niobium and tantalum exist as volatile solids or monomeric trigonal bipyramidal molecules in the vapour phase. Thus, DbCl5 is expected to be a volatile solid. Similarly, the pentafluoride, DbF5, should be even more volatile. Hydrolysis of the halides is known to readily form the oxyhalides, MOX3. Thus the halides DbX5 should react with water to form DbOX3. The reaction with fluoride ion is also well known for the lighter homologues and dubnium is expected to form a range of fluoro-complexes. In particular, reaction of the pentafluoride with HF should form a hexafluorodubnate ion, DbF6–. Excess fluoride should lead to DbF72– and DbOF52–. If eka-tantalum properties are portrayed, higher concentrations of fluoride should ultimately form DbF83– since NbF83– is not known.
Dubnium
5
Experimental chemistry Gas phase chemistry The chemistry of dubnium has been studied for several years using gas thermochromatography. The experiments have studied the relative adsorption characteristics of isotopes of niobium, tantalum and dubnium radioisotopes. The results have indicated the formation of typical group 5 halides and oxyhalides, namely DbCl5, DbBr5, DbOCl3 and DbOBr3. Reports on these early experiments usually refer to dubnium as hahnium.
Summary of compounds and complex ions Formula
Names(s)
DbCl5
dubnium pentachloride ; dubnium(V) chloride
DbBr5
dubnium pentabromide ; dubnium(V) bromide
DbOCl3
dubnium oxychloride ; dubnium(V) trichloride oxide ; dubnyl(V) chloride
DbOBr3
dubnium oxybromide ; dubnium(V) tribromide oxide ; dubnyl(V) bromide
History of synthesis of isotopes by cold fusion This section deals with the synthesis of nuclei of dubnium by so-called "cold" fusion reactions. These are processes which create compound nuclei at low excitation energy (~10-20 MeV, hence "cold"), leading to a higher probability of survival from fission. The excited nucleus then decays to the ground state via the emission of one or two neutrons only. 209
Bi(50Ti,xn)259-xDb (x=1,2,3)
The first attempts to synthesise element 105 using cold fusion reactions were performed in 1976 by the team at FLNR, Dubna using the above reaction. They were able to detect a 5 s spontaneous fission (SF) activity which they assigned to 257105. This assignment was later corrected to 258105. In 1981, the team at GSI studied this reaction using the improved technique of correlation of genetic parent-daughter decays. They were able to positively identify 258Db, the product from the 1n neutron evaporation channel.[5] In 1983, the team at Dubna revisited the reaction using the method of identification of a descendant using chemical separation. They succeeded in measuring alpha decays from known descendants of the decay chain beginning with 258105. This was taken as providing some evidence for the formation of element 105 nuclei. The team at GSI revisited the reaction in 1985 and were able to detect 10 atoms of 257Db.[6] After a significant upgrade of their facilities in 1993, in 2000 the team measured 120 decays of 257Db, 16 decays of 256Db and decay of 258 Db in the measurement of the 1n, 2n and 3n excitation functions. The data gathered for 257 Db allowed a first spectroscopic study of this isotope and identified an isomer, 257mDb, and a first determination of a decay level structure for 257Db.[7] The reaction was used in spectroscopic studies of isotopes of mendelevium and einsteinium in 2003-2004.[8]
Dubnium 209
Bi(49Ti,xn)258-xDb (x=2?)
This reaction was studied by Yuri Oganessian and the team at Dubna in 1983. They observed a 2.6 s SF activity tentatively assigned to 256Db. Later results suggest a possible reassignment to 256Rf, resulting from the ~30% EC branch in 256Db. 209
Bi(48Ti,xn)257-xDb (x=1?)
This reaction was studied by Yuri Oganessian and the team at Dubna in 1983. They observed a 1.6 s activity with a ~80% alpha branch with a ~20% SF branch. The activity was tentatively assigned to 255Db. Later results suggest a reassignment to 256Db. 208
Pb(51V,xn)259-xDb (x=1,2)
The team at Dubna also studied this reaction in 1976 and were again able to detect the 5 s SF activity, first tentatively assigned to 257Db and later to 258Db. In 2006, the team at LBNL reinvestigated this reaction as part of their odd-Z projectile program. They were able to detect 258Db and 257Db in their measurement of the 1n and 2n neutron evaporation channels. [9] 207
Pb(51V,xn)258-xDb
The team at Dubna also studied this reaction in 1976 but this time they were unable to detect the 5 s SF activity, first tentatively assigned to 257Db and later to 258Db. Instead, they were able to measure a 1.5 s SF activity, tentatively assigned to 255Db. 205
Tl(54Cr,xn)259-xDb (x=1?)
The team at Dubna also studied this reaction in 1976 and were again able to detect the 5 s SF activity, first tentatively assigned to 257Db and later to 258Db.
History of synthesis of isotopes by hot fusion This section deals with the synthesis of nuclei of dubnium by so-called "hot" fusion reactions. These are processes which create compound nuclei at high excitation energy (~40-50 MeV, hence "hot"), leading to a reduced probability of survival from fission and quasi-fission. The excited nucleus then decays to the ground state via the emission of 3-5 neutrons. 232
Th(31P,xn)263-xDb (x=5)
There are very limited reports that this rare reaction using a P-31 beam was studied in 1989 by Andreyev et al. at the FLNR. One source suggests that no atoms were detected whilst a better source from the Russians themselves indicates that 258Db was synthesised in the 5n channel with a yield of 120 pb.
6
Dubnium 238
U(27Al,xn)265-xDb (x=4,5)
In 2006, as part of their study of the use of uranium targets in superheavy element synthesis, the LBNL team led by Ken Gregorich studied the excitation functions for the 4n and 5n channels in this new reaction.[10] 236
U(27Al,xn)263-xDb (x=5,6)
This reaction was first studied by Andreyev et al. at the FLNR, Dubna in 1992. They were able to observe 258Db and 257Db in the 5n and 6n exit channels with yields of 450 pb and 75 pb, respectively.[11] 243
Am(22Ne,xn)265-xDb (x=5)
The first attempts to synthesis element 105 were performed in 1968 by the team at the Flerov Laboratory of Nuclear Reactions (FLNR) in Dubna,Russia. They observed two alpha lines which they tentatively assigned to 261105 and 260105. They repeated their experiment in 1970 looking for spontaneous fission. They found a 2.2 s SF activity which they assigned to 261105. In 1970, the Dubna team began work on using gradient thermochromatography in order to detect element 105 in chemical experiments as a volatile chloride. In their first run they detected a volatile SF activity with similar adsorption properties to NbCl5 and unlike HfCl4. This was taken to indicate the formation of nuclei of dvi-niobium as [105]Cl5. In 1971, they repeated the chemistry experiment using higher sensitivity and observed alpha decays from an dvi-niobium component, taken to confirm the formation of 260105. The method was repeated in 1976 using the formation of bromides and obtained almost identical results, indicating the formation of a volatile, dvi-niobium-like [105]Br5. 241
Am(22Ne,xn)263-xDb (x=4,5)
In 2000, Chinese scientists at the Institute of Modern Physics (IMP), Lanzhou, announced the discovery of the previously unknown isotope 259Db formed in the 4n neutron evaporation channel. They were also able to confirm the decay properties for 258Db.[12] 248
Cm(19F,xn)267-xDb (x=4,5)
This reaction was first studied in 1999 at the Paul Scherrer Institute (PSI) in order to produce 262Db for chemical studies. Just 4 atoms were detected with a cross section of 260 pb.[13] Japanese scientists at JAERI studied the reaction further in 2002 and determined yields for the isotope 262Db during their efforts to study the aqueous chemistry of dubnium.[14] 249
Bk(18O,xn)267-xDb (x=4,5)
Following from the discovery of 260Db by Albert Ghiorso in 1970 at the University of California (UC), the same team continued in 1971 with the discovery of the new isotope 262 Db. They also observed an unassigned 25 s SF activity, probably associated with the now-known SF branch of 263Db.[15] In 1990, a team led by Kratz at LBNL definitively discovered the new isotope 263Db in the 4n neutron evaporation channel.[16] This reaction has been used by the same team on several occasions in order to attempt to confirm an electron capture (EC) branch in 263Db leading to long-lived 263Rf (see rutherfordium).[17]
7
Dubnium
8
249
Bk(16O,xn)265-xDb (x=4)
Following from the discovery of 260Db by Albert Ghiorso in 1970 at the University of California (UC), the same team continued in 1971 with the discovery of the new isotope 261 Db.[15] 250
Cf(15N,xn)265-xDb (x=4)
Following from the discovery of 260Db by Ghiorso in 1970 at LBNL, the same team continued in 1971 with the discovery of the new isotope 261Db.[15] 249
Cf(15N,xn)264-xDb (x=4)
In 1970, the team at the Lawrence Berkeley National Laboratory (LBNL) studied this reaction and identified the isotope 260105 in their discovery experiment. They used the modern technique of correlation of genetic parent-daughter decays to confirm their assignment.[18] In 1977, the team at Oak Ridge repeated the experiment and were able to confirm the discovery by the identification of K X-rays from the daughter lawrencium.[19] 254
Es(13C,xn)267-xDb
In 1988, scientists as the Lawrence Livermore National Laboratory (LLNL) used the asymmetric hot fusion reaction with an einsteinium-254 target to search for the new nuclides 264Db and 263Db. Due to the low sensitivity of the experiment caused by the small Es-254 target,they were unable to detect any evaporation residues (ER).
Synthesis of isotopes as decay products Isotopes of dubnium have also been identified in the decay of heavier elements. Observations to date are summarised in the table below: Evaporation Residue
Observed Db isotope
288
268
287
267
282
266
267
263
115
Db
115
Db
113
Db
Bh
278
Db
113 ,
266
Bh
262
Db
265
261
272
260
Bh
Db
Rg
266
Mt ,
Db
262
Bh
258
Db
261
257
260
256
Bh Bh
Db Db
Dubnium
9
Chronology of isotope discovery Isotope
Year discovered
discovery reaction
256
1983? , 2000
209
257
1985
209
257
2000
209
258
1976? , 1981
209
259
2001
241
260
1970
249
261
1971
249
262
1971
249
263
1971? , 1990
249
264
unknown
265
unknown
266
2006
237
267
2003
243
268
2003
243
Db Dbg Dbm Db Db Db Db Db Db Db Db Db Db Db
Bi(50Ti,3n) Bi(50Ti,2n) Bi(50Ti,2n) Bi(50Ti,n) Am(22Ne,4n) Cf(15N,4n) Bk(16O,4n) Bk(18O,5n) Bk(18O,4n)
Np(48Ca,3n) Am(48CaCa,4n) Am(48Ca,3n)
Isomerism in dubnium nuclides 260
Db
Recent data on the decay of 272Rg has revealed that some decay chains continue through 260 Db with extraordinary longer life-times than expected. These decays have been linked to an isomeric level decaying by alpha decay with a half-life of ~19 s. Further research is required to allow a definite assignment. 258
Db
Evidence for an isomeric state in 258Db has been gathered from the study of the decay of 266 Mt and 262Bh. It has been noted that those decays assigned to an electron capture (EC) branch has a significantly different half-life to those decaying by alpha emission. This has been taken to suggest the existence of an isomeric state decaying by EC with a half-life of ~20 s. Further experiments are required to confirm this assignment. 257
Db
A study of the formation and decay of 257Db has proved the existence of an isomeric state. Initially, 257Db was taken to decay by alpha emission with energies 9.16,9.07 and 8.97 MeV. A measurement of the correlations of these decays with those of 253Lr have shown that the 9.16 MeV decay belongs to a separate isomer. Analysis of the data in conjunction with theory have assigned this activity to a meta stable state, 257mDb. The ground state decays
Dubnium
10
by alpha emission with energies 9.07 and 8.97 MeV. Spontaneous fission of not confirmed in recent experiments.
257m,g
Db was
Spectroscopic decay level schemes for dubnium isotopes 257
Db
Retracted isotopes 255
Db
In 1983, scientists at Dubna carried out a series of supportive experiments in their quest for the discovery of element 107. In two such experiments, they claimed they had detected a ~1.5 s spontaneous fission activity from the reactions 207 51 Pb( V,xn) and 209 48 Bi( Ti,xn). The activity was assigned to 255Db. Later research suggested that the assignment should be changed to 256Db. As such, the isotope 255 Db is currently not This is the currently suggested decay level scheme for 257Dbg,m from recognised on the chart of the study performed in 2001 by Hessberger et al. at GSI radionuclides and further research is required to confirm this isotope.
Chemical yields of isotopes Cold fusion The table below provides cross-sections and excitation energies for cold fusion reactions producing dubnium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
CN
1n
2n
51
208
259
1.54 nb , 15.6 MeV
1.8 nb , 23.7 MeV
50
209
259
4.64 nb , 16.4 MeV
2.4 nb , 22.3 MeV
V Ti
Pb Bi
Db Db
3n
200 pb , 31.0 MeV
Dubnium
11
Hot fusion The table below provides cross-sections and excitation energies for hot fusion reactions producing dubnium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
CN
3n
4n
5n
27
238
265
+
+
22
241
263
1.6 nb
3.6 nb
22
243
265
+
+
19
248
267
18
249
267
Al Ne Ne F O
U Am Am Cm Bk
Db Db Db
1.0 nb
Db Db
10.0 nb
6.0 nb
External links • WebElements.com - Dubnium
[20]
References [1] Data from XU_187225_1.ens (http:/ / www. nndc. bnl. gov/ useroutput/ XU_187225_1. html) [2] " TWG report (http:/ / www. iupac. org/ publications/ pac/ 1993/ pdf/ 6508x1757. pdf)". . Retrieved 2009-05-05. [3] " IUPAC 1994 recomm (http:/ / www. iupac. org/ publications/ pac/ 1994/ pdf/ 6612x2419. pdf)". . Retrieved 2009-05-06. [4] " IUPAC 1997 recomm (http:/ / www. iupac. org/ publications/ pac/ 1997/ pdf/ 6912x2471. pdf)". . Retrieved 2009-05-05. [5] Munzenberg et al. (1981). " Identification of element 107 by α correlation chains (http:/ / www. springerlink. com/ content/ tlx33361417u3n26/ )". Z. Phys. A. 300: 1. doi: 10.1007/BF01412623 (http:/ / dx. doi. org/ 10. 1007/ BF01412623). . [6] Hessberger et al. (1985). " The new isotopes 258105,257105,254Lr and 253Lr" (http:/ / www. springerlink. com/ content/ j04846712485p483/ )". Z. Phys A. 322: 4. doi: 10.1007/BF01415134 (http:/ / dx. doi. org/ 10. 1007/ BF01415134). . [7] F. P. Hessberger et al. (2001). " Decay properties of neutron-deficient isotopes 256,257Db, 255Rf, 252,253Lr (http:/ / www. edpsciences. org/ articles/ epja/ abs/ 2001/ 09/ epja1103/ epja1103. html)". Eur. Phys. J. A 12: 57-67. . [8] F. P. Hessberger et al. (2005). Eur. Phys. J A. 26: 2. doi: Energy systematics of low-lying Nilsson levels in odd-mass einsteinium isotopes (http:/ / dx. doi. org/ Energy+ systematics+ of+ low-lying+ Nilsson+ levels+ in+ odd-mass+ einsteinium+ isotopes). http:/ / www. springerlink. com/ content/ 7n66l7650112m776/ . [9] Gates et al. (2005). " Measurement of the 208Pb(51V, xn)259-xDb Excitation Function (http:/ / rnc. lbl. gov/ nsd/ annualreport2005/ contributions/ Gates1_LE. pdf)". LBNL Annual Report. . [10] " 238U studies (http:/ / rnc. lbl. gov/ nsd/ annualreport2005/ contributions/ Gregorich_LE. pdf)". . Retrieved 2009-05-05. [11] A. N. Andreyev et al. (1992). " Investigation of the fusion reaction 27Al + 236U → 263105 at excitation energies of 57 MeV and 65 MeV (http:/ / www. springerlink. com/ content/ p47303w3887q36r7/ )". Z. Phys. A. 344: 2. doi: 10.1007/BF01291709 (http:/ / dx. doi. org/ 10. 1007/ BF01291709). . [12] Z. G. Gan et al. (2001). " A new alpha-particle-emitting isotope 259Db (http:/ / www. springerlink. com/ content/ gflp9qfdqp0l9r1u/ )". Eur. Phys. J. A 10: 1. doi: 10.1007/s100500170140 (http:/ / dx. doi. org/ 10. 1007/ s100500170140). . [13] R. Dressler et al. (1999). " Production of 262Db (Z=105) in the reaction 248Cm(19F, 5n) (http:/ / prola. aps. org/ abstract/ PRC/ v59/ i6/ p3433_1)". Phys. Rev. C 59: 3433 - 3436. doi: 10.1103/PhysRevC.59.3433 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 59. 3433). . [14] Y. Nagame et al. (2002). " Production Cross Sections of 261Rf and 262Db in Bombardments of 248Cm with 18O and 19F Ions (http:/ / sciencelinks. jp/ j-east/ article/ 200219/ 000020021902A0652005. php)". J. Nucl. Radiochem. Sci 3: 85-88. .
Dubnium [15] A. Ghiorso et al. (1971). " Two New Alpha-Particle Emitting Isotopes of Element 105, 261Ha and 262Ha (http:/ / prola. aps. org/ abstract/ PRC/ v4/ i5/ p1850_1)". Phys. Rev. C 4: 1850 - 1855. doi: 10.1103/PhysRevC.4.1850 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 4. 1850). . [16] J. V. Kratz et al. (1992). " New nuclide 263Ha (http:/ / prola. aps. org/ abstract/ PRC/ v45/ i3/ p1064_1)". Phys. Rev. C 45: 1064 - 1069. doi: 10.1103/PhysRevC.45.1064 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 45. 1064). . [17] " EC of 263Db (http:/ / www. ulrich-rieth. de/ publikationen/ RCA0301_059. PDF)". . Retrieved 2009-05-05. [18] A. Ghiorso et al. (1970). " New Element Hahnium, Atomic Number 105 (http:/ / prola. aps. org/ abstract/ PRL/ v24/ i26/ p1498_1)". Phys. Rev. Lett. 24: 1498 - 1503. doi: 10.1103/PhysRevLett.24.1498 (http:/ / dx. doi. org/ 10. 1103/ PhysRevLett. 24. 1498). . [19] C. E. Bemis et al. (1977). " Production, L x-ray identification, and decay of the nuclide 260105 (http:/ / prola. aps. org/ abstract/ PRC/ v16/ i3/ p1146_1)". Phys. Rev. C 16: 1146 - 1158. . [20] http:/ / www. webelements. com/ webelements/ elements/ text/ Db/ index. html
12
Article Sources and Contributors
Article Sources and Contributors Dubnium Source: http://en.wikipedia.org/w/index.php?oldid=306272502 Contributors: Adashiel, Aesopos, Ahoerstemeier, Alan Liefting, AlimanRuna, Andres, BlueEarth, Borislav Dopudja, Bryan Derksen, Carnildo, Closedmouth, Conversion script, Curps, D64, Darrien, DragonflySixtyseven, Drjezza, DyingIce, EPO, Edgar181, Emperorbma, Eric119, Femto, Gamma, Greatpatton, Heian-794, Heron, IanA, Icairns, Ideyal, Irregulargalaxies, J.delanoy, John, Kalamkaar, Karelj, Karlhahn, Kgf0, Kingdon, Ksbrown, Kwamikagami, Lethalgeek, LilHelpa, Marcika, MarsRover, Martin451, Materialscientist, Mav, Mdf, Mimihitam, N-true, Neil916, Nergaal, Nk, Pixel ;-), PlatinumX, Pne, Poolkris, Pras, Psyche825, Pt, Puchiko, Quercusrobur, Remember, Rich Farmbrough, Rifleman 82, Roentgenium111, Romanm, Romanskolduns, Rursus, Sakus, Saperaud, Schneelocke, Simon J Kissane, Sl, SpaceFlight89, Stifynsemons, Stone, StradivariusTV, Tagishsimon, Tetracube, Trigger820, UnitedStatesian, Until It Sleeps, Useight, Vina, Waase, Warut, Writtenright, Yekrats, 92 anonymous edits
Image Sources, Licenses and Contributors image:Db-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Db-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier Image:257Db decay scheme.png Source: http://en.wikipedia.org/w/index.php?title=File:257Db_decay_scheme.png License: Public Domain Contributors: User:Drjezza
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
13
Seaborgium
1
Seaborgium dubnium ← seaborgium → bohrium W ↑ Sg ↓ (Uph) Periodic Table Extended Periodic Table General Name, symbol, number Element category Group, Period, Block Appearance
seaborgium, Sg, 106 transition metals 6, 7, d unknown, probably silvery white or metallic gray
Standard atomic weight
[271] g·mol
Electron configuration
[Rn] 7s 5f
Electrons per shell
−1
2
14
4
6d
2, 8, 18, 32, 32, 12, 2 Physical properties
Phase
presumably a solid Density (near r.t.)
−3
unknown g·cm Atomic properties
Crystal structure Oxidation states
unknown 6
Atomic radius (calc.)
unknown pm
Covalent radius
unknown pm Miscellaneous
CAS registry number
54038-81-2 Most-stable isotopes
Seaborgium
2
Main article: Isotopes of seaborgium iso 271
Sg
NA syn
half-life 1.9 min
DM 67% α
DE (MeV)
DP
8.54
267
8.20
263
Rf
33% SF 267
Sg
syn
1.4 min
17% α
Rf
83% SF syn
0.36 s
SF
265b
Sg
syn
16.2 s
α
8.70
261g
265a
Sg
syn
8.9 s
α
8.90,8.84,8.76
261
264
syn
68 ms
SF
syn
0.9 s
87% α
9.25
259
9.06
259
9.62,9.55,9.47,9.42,9.37
257g
266
Sg
Sg
263m
Sg
Rf
Rf
Rf
13% SF 263g
syn
0.3 s
262
syn
15 ms
SF
261
syn
0.18 s
98.1% α
Sg
Sg Sg
α
1.3% ε
Rf
Rf
261
Db
0.6% SF 260
Sg
syn
3.6 ms
26% α
9.81,9.77,9.72
256
9.62,9.36,9.03
255
Rf
74% SF 259
syn
0.48 s
α
258
syn
2.9 ms
SF
Sg Sg
Rf
References
Seaborgium (pronounced /siːˈbɔrɡiəm/) is a chemical element in the periodic table that has the symbol Sg and atomic number 106. Seaborgium is a synthetic element whose most stable isotope 271Sg has a half-life of 1.9 minutes. Chemistry experiments with seaborgium have firmly placed it in group 6 as a heavier homologue to tungsten.
Proposed names (main article: Element naming controversy) The Berkeley team suggested the name seaborgium (Sg) to honor the American chemist Glenn T. Seaborg credited as a member of the American group in recognition of his participation in the discovery of several other actinides. The name selected by the team became controversial (see element naming controversy). The IUPAC adopted unnilhexium (symbol Unh) as a temporary, systematic element name. In 1994 a committee of IUPAC
Seaborgium
3
recommended that element 106 be named rutherfordium and adopted a rule that no element can be named after a living person.[1] This ruling was fiercely objected to by the American Chemical Society. Critics pointed out that a precedent had been set in the naming of einsteinium during Albert Einstein's life and a survey indicated that chemists were not concerned with the fact that Seaborg was still alive. In 1997, as part of a compromise involving elements 104 to 108, the name seaborgium for element 106 was recognized internationally.[2]
Electronic structure Seaborgium is element 106 in the Periodic Table. The two forms of the projected electronic structure are: Bohr model Quantum mechanical model
2, 8, 18, 32, 32, 12, 2 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f146d4
Extrapolated chemical properties of eka-tungsten/dvi-molybdenum Oxidation states Element 106 is projected to be the third member of the 6d series of transition metals and the heaviest member of group 6 in the Periodic Table, below chromium, molybdenum and tungsten. All the members of the group readily portray their group oxidation state of +6 and the state becomes more stable as the group is descended. Thus seaborgium is expected to form a stable +6 state. For this group, stable +5 and +4 states are well represented for the heavier members and the +3 state is known but reducing, except for chromium(III).
Chemistry Much seaborgium chemical behavior is predicted by extrapolation from its lighter cogeners molybdenum and tungsten. Molybdenum and tungsten readily form stable trioxides MO3, so seaborgium should form SgO3. The oxides MO3 are soluble in alkali with the formation of oxyanions, so seaborgium should form a seaborgate ion, SgO42−. In addition, WO3 reacts with acid, suggesting similar amphotericity for SgO3. Molybdenum oxide, MoO3, also reacts with moisture to form a hydroxide MoO2(OH)2, so SgO2(OH)2 is also feasible. The heavier homologues readily form the volatile, reactive hexahalides MX6 (X=Cl,F). Only tungsten forms the unstable hexabromide, WBr6. Therefore, the compounds SgF6 and SgCl6 are predicted, and "eka-tungsten character" may show itself in increased stability of the hexabromide, SgBr6. These halides are unstable to oxygen and moisture and readily form volatile oxyhalides, MOX4 and MO2X2. Therefore SgOX4 (X=F,Cl) and SgO2X2 (X=F,Cl) should be possible. In aqueous solution, a variety of anionic oxyfluoro-complexes are formed with fluoride ion, examples being MOF5− and MO3F33−. Similar seaborgium complexes are expected.
Seaborgium
4
Experimental chemistry Gas phase chemistry Initial experiments aiming at probing the chemistry of seaborgium focused on the gas thermochromatography of a volatile oxychloride. Seaborgium atoms were produced in the reaction 248Cm(22Ne,4n)266Sg, thermalised, and reacted with an O2/HCl mixture. The adsorption properties of the resulting oxychloride were measured and compared with those of molybdenum and tungsten compounds. The results indicated that seaborgium formed a volatile oxychloride akin to those of the other group 6 elements: Sg + O2 + 2 HCl → SgO2Cl2 + H2 In 2001, a team continued the study of the gas phase chemistry of seaborgium by reacting the element with O2 in a H2O environment. In a manner similar to the formation of the oxychloride, the results of the experiment indicated the formation of seaborgium oxide hydroxide, a reaction well known among the lighter group 6 homologues.[3] 2 Sg + 3 O2 → 2 SgO3 SgO3 + H2O → SgO2(OH)2
Aqueous phase chemistry In its aqueous chemistry, seaborgium has been shown to resemble its lighter homologues molybdenum and tungsten, forming a stable +6 oxidation state. Seaborgium was eluted from cation exchange resin using a HNO3/HF solution, most likely as neutral SgO2F2 or the anionic complex ion [SgO2F3]−. In contrast, in 0.1 M HNO3, seaborgium does not elute, unlike Mo and W, indicating that the hydrolysis of [Sg(H2O)6]6+ only proceeds as far as the cationic complex [Sg(OH)5(H2O)]+.
Summary of investigated compounds and complex ions Formula
Names(s)
SgO2Cl2
seaborgium oxychloride ; seaborgium(VI) dioxide dichloride ; seaborgyl dichloride
SgO2F2
seaborgium oxyfluoride ; seaborgium(VI) dioxide difluoride ; seaborgyl difluoride
SgO3
seaborgium oxide ; seaborgium(VI) oxide ; seaborgium trioxide
SgO2(OH)2
seaborgium oxide hydroxide ; seaborgium(VI) dioxide dihydroxide
[SgO2F3]−
trifluorodioxoseaborgate(VI)
[Sg(OH)5(H2O)]+
aquapentahydroxyseaborgium(VI)
Seaborgium
History of synthesis of isotopes by cold fusion This section deals with the synthesis of nuclei of seaborgium by so-called "cold" fusion reactions. These are processes which create compound nuclei at low excitation energy (~10-20 MeV, hence "cold"), leading to a higher probability of survival from fission. The excited nucleus then decays to the ground state via the emission of one or two neutrons only. 208
Pb(54Cr,xn)262-xSg (x=1,2,3)
The first attempt to synthesise element 106 in cold fusion reactions was performed in September 1974 by a Soviet team led by G. N. Flerov at the Joint Institute for Nuclear Research at Dubna. They reported producing a 0.48 s spontaneous fission (SF) activity which they assigned to the isotope 259106. Based on later evidence it was suggested that the team most likely measured the decay of 260Sg and its daughter 256Rf. The TWG concluded that, at the time, the results were insufficiently convincing.[4] The Dubna team revisited this problem in 1983-1984 and were able to detect a 5 ms SF activity assigned directly to 260Sg.[4] The team at GSI studied this reaction for the first time in 1985 using the improved method of correlation of genetic parent-daughter decays. They were able to detect 261Sg (x=1) and 260 Sg and measured a partial 1n neutron evaporation excitation function. [5] In December 2000, the reaction was studied by a team at GANIL, France and were able to detect 10 atoms of 261Sg and 2 atoms of 260Sg to add to previous data on the reaction. After a facility upgrade, the GSI team measured the 1n excitation function in 2003 using a metallic lead target. Of significance, in May 2003, the team successfully replaced the lead-208 target with more resistant lead(II) sulfide targets (PbS) which will allow more intense beams to be used in the future. They were able to measure the 1n,2n and 3n excitation functions and performed the first detailed alpha-gamma spectroscopy on the isotope 261Sg. They detected ~1600 atoms of the isotope and identified new alpha lines as well as measuring a more accurate half-life and new EC and SF branchings. Furthermore, they were able to detect the K X-rays from the daughter rutherfordium element for the first time. They were also able to provide improved data for 260Sg, including the tentative observation of an isomeric level. The study was continued in September 2005 and March 2006. The accumulated work on 261Sg was published in 2007. [6] Work in September 2005 also aimed to begin spectroscopic studies on 260Sg. 207
Pb(54Cr,xn)261-xSg (x=1,2)
The team at Dubna also studied this reaction in 1974 with identical results as for their first experiments with a Pb-208 target. The SF activities were first assigned to 259Sg and later to 260 Sg and/or 256Rf. Further work in 1983-1984 also detected a 5 ms SF activity assigned to the parent 260Sg.[4] The GSI team studied this reaction for the first time in 1985 using the method of correlation of genetic parent-daughter decays. They were able to positively identify 259Sg as a product from the 2n neutron evaporation channel.[5] The reaction was further used in March 2005 using PbS targets to begin a spectroscopic study of the even-even isotope 260Sg.
5
Seaborgium 206
Pb(54Cr,xn)260-xSg
This reaction was studied in 1974 by the team at Dubna. It was used to assist them in their assignment of the observed SF activities in reactions using Pb-207 and Pb-208 targets. They were unable to detect any SF, indicating the formation of isotopes decaying primarily by alpha decay.[4] 208
Pb(52Cr,xn)260-xSg (x=1,2)
The team at Dubna also studied this reaction in their series of cold fusion reactions performed in 1974. Once again they were unable to detect any SF activities.[4] The reaction was revisited in 2006 by the team at LBNL as part of their studies on the effect of the isospin of the projectile and hence the mass number of the compound nucleus on the yield of evaporation residues. They were able to identify 259Sg and 258Sg in their measurement of the 1n excitation function.[7] 209
Bi(51V,xn)260-xSg (x=2)
The team at Dubna also studied this reaction in their series of cold fusion reactions performed in 1974. Once again they were unable to detect any SF activities.[4] In 1994, the synthesis of seaborgium was revisited using this reaction by the GSI team, in order to study the new even-even isotope 258Sg. Ten atoms of 258Sg were detected and decayed by spontaneous fission.
History of synthesis of isotopes by hot fusion This section deals with the synthesis of nuclei of seaborgium by so-called "hot" fusion reactions. These are processes which create compound nuclei at high excitation energy (~40-50 MeV, hence "hot"), leading to a reduced probability of survival from fission and quasi-fission. The excited nucleus then decays to the ground state via the emission of 3-5 neutrons. 238
U(30Si,xn)268-xSg (x=3,4,5,6)
This reaction was first studied by Japanese scientists at the Japan Atomic Energy Research Institute (JAERI) in 1998. They detected a spontaneous fission activity which they tentatively assigned to the new isotope 264Sg or 263Db, formed by EC of 263Sg.[8] In 2006, the teams at GSI and LBNL both studied this reaction using the method of correlation of genetic parent-daughter decays. The LBNL team measured an excitation function for the 4n,5n and 6n channels, whilst the GSI team were able to observe an additional 3n activity.[9] [10] [11] Both teams were able to identify the new isotope 264Sg which decayed with a short lifetime by spontaneous fission. 248
Cm(22Ne,xn)270-xSg (x=4?,5)
In 1993, at Dubna, Yuri Lazarev and his team announced the discovery of long-lived 266Sg and 265Sg produced in the 4n and 5n channels of this nuclear reaction following the search for seaborgium isotopes suitable for a first chemical study. It was announced that 266Sg decayed by 8.57 MeV alpha-particle emission with a projected half-life of ~20 s, lending strong support to the stabilising effect of the Z=108,N=162 closed shells.[12] This reaction was studied further in 1997 by a team at GSI and the yield, decay mode and half-lives for
6
Seaborgium
7
266
Sg and 265Sg have been confirmed, although there are still some discrepancies. In the recent synthesis of 270Hs (see hassium), 266Sg was found to undergo exclusively SF with a short half-life (TSF = 360 ms). It is possible that this is the ground state, (266gSg) and that the other activity, produced directly, belongs to a high spin K-isomer, 266mSg, but further results are required to confirm this. A recent re-evaluation of the decay characteristics of 265Sg and 266Sg has suggested that all decays to date in this reaction were in fact from 265Sg, which exists in two isomeric forms. The first, 265aSg has a principal alpha-line at 8.85 MeV and a calculated half-life of 8.9 s, whilst 265bSg has a decay energy of 8.70 MeV and a half-life of 16.2 s. Both isomeric levels are populated when produced directly. Data from the decay of 269Hs indicates that 265b Sg is produced during the decay of 269Hs and that 265bSg decays into the shorter-lived 261g Rf isotope. This means that the observation of 266Sg as a long-lived alpha emitter is retracted and that it does indeed undergo fission in a short time. Regardless of these assignments, the reaction has been successfully used in the recent attempts to study the chemistry of seaborgium (see below). 249
Cf(18O,xn)267-xSg (x=4)
The synthesis of element 106 was first attempted in 1974 by the team at LBNL. In their discovery experiment, they were able to apply the new method of correlation of genetic parent-daughter decays to identify the new isotope 263Sg. In 1975, the team at Oak Ridge were able to confirm the decay data but were unable to identify coincident X-rays in order to prove that seaborgium as produced. In 1979, the team at Dubna studied the reaction by detection of SF activities. In comparison with data from Berkeley, they calculated a 70% SF branching for 263Sg. The synthesis and discovery reaction was confirmed in 1994 by a different team at LBNL. [13]
Synthesis of isotopes as decay products Isotopes of seaborgium have also been observed in the decay of heavier elements. Observations to date are summarised in the table below: Evaporation Residue 291
116 ,
287
114 ,
283
112
Observed Sg isotope 271
Sg
271
267
270
266
Hs
Sg
Hs
277
Sg
Uub ,
271
Ds ,
273
Ds ,
267
Ds
270
Ds ,
264
Hs
Hs
265
Sg
263
Sg
262
Ds
269
269
Sg
265
Hs
261
Sg
260
Sg
Seaborgium
8
Chronology of isotope discovery Isotope
Year discovered
discovery reaction
258
1994
209
259
1985
207
260
1985
208
261
1985
208
262
2001
207
263
1974
249
263
1994
208
264
2006
238
265
1993
248
266
2004
248
267
2004
248
268
unknown
269
unknown
270
unknown
271
2003
Sg Sg Sg Sg Sg Sgm Sgg Sg Sg Sg Sg Sg Sg Sg Sg
Bi(51V,2n) Pb(54Cr,2n) Pb(54Cr,2n) Pb(54Cr,n) Pb(64Ni,n)
[14]
Cf(18O,4n) Pb(64Ni,n)
[14]
U(30Si,4n) Cm(22Ne,5n) Cm(26Mg,4n) Cm(26Mg,3n)
242
Pu(48Ca,3n)
[15]
[16]
Isotopes There are 11 known isotopes of seaborgium (excluding meta-stable and K-spin isomers). The longest-lived is 271Sg which decays through alpha decay and spontaneous fission. It has a half-life of 1.9 minutes. The shortest-lived isotope is 258Sg which also decays through alpha decay and spontaneous fission. It has a half-life of 2.9 ms.
Isomerism in seaborgium nuclides 266
Sg
Initial work identified an 8.63 MeV alpha-decaying activity with a half-life of ~21s and assigned to the ground state of 266Sg. Later work identified a nuclide decaying by 8.52 and 8.77 MeV alpha emission with a half-life of ~21s, which is unusual for an even-even nuclide. Recent work on the synthesis of 270Hs identified 266Sg decaying by SF with a short 360 ms half-life. The recent work on 277112 and 269Hs has provided new information on the decay of 265Sg and 261Rf. This work suggested that the initial 8.77 MeV activity should be reassigned to 265Sg. Therefore the current information suggests that the SF activity is the ground state and the 8.52 MeV activity is a high spin K-isomer. Further work is required to confirm these assignments. A recent re-evaluation of the data has suggested that the 8.52 MeV activity should be associated with 265Sg and that 266Sg only undergoes fission.
Seaborgium 265
Sg
The recent direct synthesis of 265Sg resulted in four alpha-lines at 8.94,8.84,8.76 and 8.69 MeV with a half-life of 7.4 seconds. The observation of the decay of 265Sg from the decay of 277 112 and 269Hs indicated that the 8.69 MeV line may be associated with an isomeric level with an associated half-life of ~ 20 s. It is plausible that this level is causing confusion between assignments of 266Sg and 265Sg since both can decay to fissioning rutherfordium isotopes. A recent re-evaluation of the data has indicated that there are indeed two isomers, one with a principal decay energy of 8.85 MeV with a half-life of 8.9 s, and a second isomer which decays with energy 8.70 MeV with a half-life of 16.2 s. 263
Sg
The discovery synthesis of 263Sg resulted in an alpha-line at 9.06 MeV. Observation of this nuclide by decay of 271gDs, 271mDs and 267Hs has confirmed an isomer decaying by 9.25 MeV alpha emission. The 9.06 MeV decay was also confirmed. The 9.06 MeV activity has been assigned to the ground state isomer with an associated half-life of 0.3 s. The 9.25 MeV activity has been assigned to an isomeric level decaying with a half-life of 0.9 s. Recent work on the synthesis of 271g,mDs was resulted in some confusing data regarding the decay of 267Hs. In one such decay, 267Hs decayed to 263Sg which decayed by alpha emission with a half-life of ~ 6 s. This activity has not yet been positively assigned to an isomer and further research is required.
9
Seaborgium
10
Spectroscopic decay schemes for seaborgium isotopes 261
Sg
Retracted isotopes 269
Sg
In the claimed synthesis of 293 118 in 1999 the isotope 269 Sg was identified as a daughter product. It decayed by 8.74 MeV alpha emission with a half-life of 22 s. The claim was retracted in 2001 and thus this seaborgium isotope is currently unknown or unconfirmed.[17]
Chemical yields of isotopes Cold fusion The table below provides This is the currently accepted decay scheme for 261Sg from the study by Streicher et al. at GSI in 2003-2006 cross-sections and excitation energies for cold fusion reactions producing seaborgium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
CN
54
207
261
54
208
262
51
209
260
52
208
260
Cr Cr V Cr
Pb Pb Bi Pb
1n
2n
3n
Sg Sg
4.23 nb , 13.0 MeV
38 pb , 21.5 MeV
Sg Sg
500 pb
281 pb , 11.0 MeV
10 pb
Seaborgium
11
Hot fusion The table below provides cross-sections and excitation energies for hot fusion reactions producing seaborgium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
CN
270
~25 pb
~250 pb
267
+
268
22
248
18
249
Ne O
Cm Cf
Sg Sg Sg
+
5n ~ 80 pb , 51.0 MeV
238
U
4n 9 pb, 40.0
30
Si
3n
6n ~30 pb , 58.0 MeV
External links • WebElements.com - Seaborgium
[18]
References [1] http:/ / www. iupac. org/ publications/ pac/ 1994/ pdf/ 6612x2419. pdf (IUPAC 1994 recomm) [2] http:/ / www. iupac. org/ publications/ pac/ 1997/ pdf/ 6912x2471. pdf (IUPAC 1997 recomm) [3] "Physico-chemical characterization of seaborgium as oxide hydroxide" (http:/ / www-w2k. gsi. de/ kernchemie/ images/ pdf_Artikel/ Radiochim_Acta_89_737_2001. pdf), Huebener et al., Radiochim. Acta, 89, 737–741 (2001).Retrieved on 2008-02-29 [4] http:/ / www. iupac. org/ publications/ pac/ 1993/ pdf/ 6508x1757. pdf (TWG report) [5] "The isotopes 259106,260106, and 261106" (http:/ / www. springerlink. com/ content/ kw84567751181537/ ), Munzenberg et al., Z. Phys. A, 1985, 322, 2.Retrieved on 2008-02-29 [6] "Alpha-Gamma Decay Studies of 261Sg" (http:/ / adsabs. harvard. edu/ abs/ 2007AcPPB. . 38. 1561S), Streicher et al., Acta Physica Polonica B, vol. 38, Issue 4, p.1561, 2007. Retrieved on 2008-03-04 [7] "Measurement of the 208Pb(52Cr,n)259Sg Excitation Function" (http:/ / rnc. lbl. gov/ nsd/ annualreport2005/ contributions/ Folden_LE. pdf), Folden et al., LBNL Annual Report 2005. Retrieved on 2008-02-29 [8] "First evidence for a new spontaneous fission decay produced in the reaction 30Si +238U" (http:/ / www. springerlink. com/ content/ wbcq11tfp6dmg486/ ), Ikezoe et al., Eur. Phys. J. A., 1998, 2, 4. Retrieved on 2008-02-29 [9] "Production of seaborgium isotopes in the reaction of 30Si + 238U" (http:/ / www. gsi. de/ informationen/ wti/ library/ scientificreport2006/ PAPERS/ NUSTAR-SHE-05. pdf), Nishio et al., GSI Annual Report 2006. Retrieved on 2008-02-29 [10] "Measurement of evaporation residue cross-sections of the reaction 30Si + 238U at subbarrier energies" (http:/ / publish. edpsciences. org/ articles/ epja/ abs/ 2006/ 12/ 10050_2006_Article_100226/ 10050_2006_Article_100226. html), Nishio et al., Eur. Phys. J. A, 29, 281-287 (2006). Retrieved on 2008-02-29 [11] "New isotope 264Sg and decay properties of 262-264Sg" (http:/ / repositories. cdlib. org/ cgi/ viewcontent. cgi?article=5547& context=lbnl), Gregorich et al., LBNL repositories. Retrieved on 2008-02-29 [12] "Discovery of Enhanced Nuclear Stability near the Deformed Shells N=162 and Z=108" (http:/ / prola. aps. org/ abstract/ PRL/ v73/ i5/ p624_1), Lazarev et al., Phys. Rev. Lett., 73, 624-627 (1994).Retrieved on 2008-02-29 [13] "First confirmation of the discovery of element 106" (http:/ / prola. aps. org/ abstract/ PRL/ v72/ i10/ p1423_1), Gregorich et al., Phys. Rev. Lett. 72, 1423-1426 (1994). Retrieved on 2008-03-04 [14] [15] [16] [17] [18]
see darmstadtium see hassium see ununquadium see ununoctium http:/ / www. webelements. com/ webelements/ elements/ text/ Sg/ index. html
Article Sources and Contributors
Article Sources and Contributors Seaborgium Source: http://en.wikipedia.org/w/index.php?oldid=306272376 Contributors: Ahoerstemeier, AlimanRuna, Andre Engels, Andres, BlueEarth, Bryan Derksen, Can't sleep, clown will eat me, Carnildo, Coldphoenix182, Conversion script, DOSGuy, DabMachine, Dante Alighieri, Darrien, Davewild, David Shay, Drjezza, EPO, Edgar181, Emperorbma, Encyclopedia77, Epbr123, Eras-mus, Eric119, Femto, Firien, Gamma, Giftlite, Glenn4pr, Greatpatton, Hadal, Hqb, Icairns, Ideyal, IvanLanin, Jeffq, Jiang, Jmundo, Joelfurr, John, Juliancolton, Jóna Þórunn, Kalamkaar, Karelj, Karl-Henner, Kgf0, Kingdon, Kurykh, Kwamikagami, LiDaobing, LilHelpa, Madame Sosostris, Matts computer, Mav, Mercury, Micothemagic, Miss Madeline, Nergaal, Nick, Nihiltres, Ohconfucius, Oncogene, Ouro, Permethius, PlatinumX, Poolkris, Psyche825, Radiojon, Railsmart, Remember, RexNL, Rifleman 82, Rjwilmsi, Roentgenium111, Romanskolduns, Salgueiro, Saperaud, Scarian, Simon J Kissane, Sl, Spongebobsquarepants, Squids and Chips, Stratocracy, Stubblyhead, Tagishsimon, Tetracube, Timc, TyrantX, Urhixidur, Vsmith, Walkerma, Warut, Yekrats, 91 anonymous edits
Image Sources, Licenses and Contributors image:Sg-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Sg-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier Image:261Sg decay scheme 2006.png Source: http://en.wikipedia.org/w/index.php?title=File:261Sg_decay_scheme_2006.png License: Public Domain Contributors: User:Drjezza
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
12
Bohrium
1
Bohrium seaborgium ← bohrium → hassium Re ↑ Bh ↓ (Ups) Periodic Table Extended Periodic Table General Name, symbol, number Element category Group, Period, Block Standard atomic weight Electron configuration Electrons per shell
bohrium, Bh, 107 transition metals 7, 7, d −1
[270] g·mol 14
[Rn] 5f
5
2
6d 7s
2, 8, 18, 32, 32, 13, 2 Physical properties
Phase
presumably a solid Atomic properties Crystal structure
Oxidation states
unknown 7 Miscellaneous
CAS registry number
54037-14-8 Most-stable isotopes
Bohrium
2
Main article: Isotopes of bohrium iso
NA
272
syn
271
syn
Bh Bh
half-life
DM
9.8 s
α
DE (MeV) 9.02
α
DP 268
Db
267
Db
270
Bh
syn
61 s
α
8.93
266
267
Bh
syn
17 s
α
8.83
263
266
syn
0.9 s
α
9.77,9.04
262
265
syn
0.9 s
α
9.24
261
264
syn
0.97 s
α
9.62,9.48
260
262m
Bh
syn
9.6 ms
α
10.37,10.24
258
262g
Bh
syn
84 ms
α
10.08,9.94,9.82,9.74,9.66
258
261
syn
11.8 ms
α
10.40,10.10,10.03
257
260
syn
35 ms
α
10.16
256
Bh Bh Bh
Bh Bh
Db Db Db Db Db Db Db Db Db
References
Bohrium (pronounced /ˈbɔəriəm/ ( listen)) is a chemical element in the periodic table that has the symbol Bh and atomic number 107. It is a synthetic element whose most stable isotope, 270Bh, has a half-life of 61 seconds. Chemical experiments confirmed bohrium's predicted position as a member of group 7 of the periodic table, as a heavier homologue to rhenium.[1]
Official discovery The first convincing synthesis was in 1981 by a German research team led by Peter Armbruster and Gottfried Münzenberg at the Gesellschaft für Schwerionenforschung (Institute for Heavy Ion Research) in Darmstadt using the Dubna reaction. 20983Bi
+
5424Cr
→
262107Bh
+n
In 1989, the GSI team successfully repeated the reaction during their efforts to measure an excitation function. During these experiments, 261Bh was also identified in the 2n evaporation channel and it was confirmed that 262Bh exists as two states - a ground state and an isomeric state. The IUPAC/IUPAP Transfermium Working Group (TWG) report in 1992 officially recognised the GSI team as discoverers of element 107.
Bohrium
3
Proposed names Historically element 107 has been referred to as eka-rhenium. The Germans suggested the name nielsbohrium with symbol Ns to honor the Danish physicist Niels Bohr. The Soviets had suggested this name be given to element 105 (which was finally called dubnium) and the German team wished to recognise both Bohr and the fact that the Dubna team had been the first to propose the cold fusion reaction. There was an element naming controversy as to what the elements from 104 to 106 were to be called; the IUPAC adopted unnilseptium (symbol Uns) as a temporary, systematic element name for this element. In 1994 a committee of IUPAC rejected the name nielsbohrium since there was no precedence for using a scientist's complete name in the naming of an element and thus recommended that element 107 be named bohrium.[2] This was opposed by the discoverers who were adamant that they had the right to name the element. The matter was handed to the Danish branch of IUPAC who voted in favour of the name bohrium. There was some concern however that the name might be confused with boron and in particular the distinguishing of the names of their respective oxo-ions bohrate and borate. Despite this, the name bohrium for element 107 was recognized internationally in 1997.[3] The IUPAC subsequently decided that bohrium salts should be called bohriates.
Electronic structure Bohrium is element 107 in the Periodic Table. The two forms of the projected electronic structure are: Bohr model Quantum mechanical model
2, 8, 18, 32, 32, 13, 2 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f146d5
Extrapolated chemical properties of eka-rhenium/dvi-technetium Oxidation states Element 107 is projected to be the fourth member of the 6d series of transition metals and the heaviest member of group VII in the Periodic Table, below manganese, technetium and rhenium. All the members of the group readily portray their group oxidation state of +7 and the state becomes more stable as the group is descended. Thus bohrium is expected to form a stable +7 state. Technetium also shows a stable +4 state whilst rhenium exhibits stable +4 and +3 states. Bohrium may therefore show these lower states as well.
Chemistry The heavier members of the group are known to form volatile heptoxides M2O7, so bohrium should also form the volatile oxide Bh2O7. The oxide should dissolve in water to form perbohric acid, HBhO4. Rhenium and technetium form a range of oxyhalides from the halogenation of the oxide. The chlorination of the oxide forms the oxychlorides MO3Cl, so BhO3Cl should be formed in this reaction. Fluorination results in MO3F and MO2F3 for the heavier elements in addition to the rhenium compounds ReOF5 and ReF7. Therefore, oxyfluoride formation for bohrium may help to indicate eka-rhenium properties.
Bohrium
4
Industrial and commercial use Like seaborgium and hassium, its neighbours, bohrium has no industrial or commercial use due to its extremely short half-life. Few atoms have ever been made, but if enough were found in one area, bohrium would constitute a radiation hazard.[4]
Experimental chemistry Gas phase chemistry In 2000, a team at the PSI conducted a chemistry reaction using atoms of 267Bh produced in the reaction between Bk-249 and Ne-22 ions. The resulting atoms were thermalised and reacted with a HCl/O2 mixture to form a volatile oxychloride. The reaction also produced isotopes of its lighter homologues, technetium (as 108Tc) and rhenium (as 169Re). The isothermal adsorption curves were measured and gave strong evidence for the formation of a volatile oxychloride with properties similar to that of rhenium oxychloride. This placed bohrium as a typical member of group 7.[1] 2 Bh + 3 O2 + 2 HCl → 2 BhO3Cl + H2
Summary of compounds Formula BhO3Cl
Names(s) bohrium oxychloride ; bohrium(VII) chloride trioxide
History of synthesis of isotopes by cold fusion This section deals with the synthesis of nuclei of bohrium by so-called "cold" fusion reactions. These are processes which create compound nuclei at low excitation energy (~10-20 MeV, hence "cold"), leading to a higher probability of survival from fission. The excited nucleus then decays to the ground state via the emission of one or two neutrons only. 209
Bi(54Cr,xn)263-xBh (x=1,2)
The synthesis of element 107 was first attempted in 1976 by scientists at the Joint Institute for Nuclear Research at Dubna using this cold fusion reaction. Analysis was by detection of spontaneous fission (SF). They discovered two SF activities, one with a 1-2 ms half-life and one with a 5 s activity. Based on the results of other cold fusion reactions, they concluded that they were due to 261107 and 257105 respectively. However, later evidence gave a much lower SF branching for 261107 reducing confidence in this assignment. The assignment of the element 105 activity was later changed to 258105, presuming that the decay of element 107 was missed. The 2 ms SF activity was assigned to 258Rf resulting from the 33% EC branch.[5] The GSI team studied the reaction in 1981 in their discovery experiments. Five atoms of 262Bh were detected using the method of correlation of genetic parent-daughter decays.[6] In 1987, an internal report from Dubna indicated that the team had been able to detect the spontaneous fission of 261107 directly. The GSI team further studied the reaction in 1989 and discovered the new isotope 261Bh during the measurement of the 1n and 2n excitation functions but were unable to detect an SF branching for 261Bh.[7] They continued their study in 2003 using newly developed bismuth(III) fluoride (BiF3) targets, used to
Bohrium provide further data on the decay data for 262Bh and the daughter 258Db. The 1n excitation function was remeasured in 2005 by the team at LBNL after some doubt about the accuracy of previous data. They observed 18 atoms of 262Bh and 3 atoms of 261Bh and confirmed the two isomers of 262Bh. [8] 209
Bi(53Cr,xn)262-xBh
The team at Dubna studied this reaction in 1976 in order to assist in their assignments of the SF activities from their experiments with a Cr-54 beam. They were unable to detect any such activity, indicating the formation of different isotopes decaying primarily by alpha decay. 209
Bi(52Cr,xn)261-xBh (x=1)
This reaction was studied for the first time in 2007 by the team at LBNL to search for the lightest bohrium isotope 260Bh. The team successfully detected 8 atoms of 260Bh decaying by correlated 10.16 MeV alpha particle emission to 256Db. The alpha decay energy indicates the continued stabilising effect of the N=152 closed shell.[9] 208
Pb(55Mn,xn)263-xBh (x=1)
The team at Dubna also studied this reaction in 1976 as part of their newly established cold fusion approach to new elements. As for the reaction using a Bi-209 target, they observed the same SF activities and assigned them to 261107 and 257105. Later evidence indicated that these should be reassigned to 258105 and 258104 (see above). In 1983, they repeated the experiment using a new technique: measurement of alpha decay from a descendant using chemical separation. The team were able to detect the alpha decay from a descendant of the 1n evaporation channel, providing some evidence for the formation of element 107 nuclei. This reaction was later studied in detail using modern techniques by the team at LBNL. In 2005 they measured 33 decays of 262Bh and 2 atoms of 261Bh, providing a 1n excitation function and some spectroscopic data of both 262Bh isomers. The 2n excitation function was further studied in a 2006 repeat of the reaction. [10] [11] The team found that this reaction had a higher 1n cross section than the corresponding reaction with a Bi-209 target, contrary to expectations. Further research is required to understand the reasons.
Synthesis of isotopes by hot fusion This section deals with the synthesis of nuclei of bohrium by so-called "hot" fusion reactions. These are processes which create compound nuclei at high excitation energy (~40-50 MeV, hence "hot"), leading to a reduced probability of survival from fission and quasi-fission. The excited nucleus then decays to the ground state via the emission of 3-5 neutrons. 238
Am(31P,xn)269-xBh (x=5?)
This reaction was first studied in 2006 at the LBNL as part of their systematic study of fusion reactions using 238U targets. Results have not been published but preliminary results appear to indicate the observation of spontaneous fission, possibly from 264Bh.[12]
5
Bohrium
6
243
Am(26Mg,xn)269-xBh (x=3,4,5)
Recently, the team at the Institute of Modern Physics (IMP), Lanzhou, have studied the nuclear reaction between americium-243 and magnesium-26 ions in order to synthesise the new isotope 265Bh [13] and gather more data on 266Bh. In two series of experiments, the team has measured partial excitation functions of the 3n,4n and 5n evaporation channels. 248
Cm(23Na,xn)271-xBh (x=5)
This reaction was studied for the first time in 2008 by the team at RIKEN, Japan, in order to study the decay properties of 266Bh, which is a decay product in their claimed decay chains of ununtrium.[14] 249
Bk(22Ne,xn)271-xBh (x=4,5)
The first attempts to synthesize element 107 by hot fusion pathways were performed in 1979 by the team at Dubna. The reaction was repeated in 1983. In both cases, they were unable to detect any spontaneous fission from nuclei of element 107. More recently, hot fusions pathways to bohrium have been re-investigated in order to allow for the synthesis of more long-lived, neutron rich isotopes to allow a first chemical study of bohrium. In 1999, the team at LBNL announced the discovery of long-lived 267Bh (5 atoms) and 266Bh (1 atom).[15] In the following year, the same team attempted to confirm the synthesis and decay of 266Bh. However, they were unable to do so and the identification of 266Bh in the first experiment is questionable. The team at the Paul Scherrer Institute (PSI) in Bern, Switzerland later synthesised 6 atoms of 267Bh in the first definitive study of the chemistry of bohrium (see below). 254
Es(16O,xn)270-xBh
As an alternative means of producing long-lived bohrium isotopes suitable for a chemical study, the synthesis of 267Bh and 266Bh were attempted in 1995 by the team at GSI using the highly asymmetric reaction using an einsteinium-254 target. They were unable to detect any product atoms.
Synthesis of isotopes as decay products Isotopes of bohrium have also been detected in the decay of heavier elements. Observations to date are shown in the table below: Evaporation Residue
Observed Bh isotope
288
272
287
271
282
270
278
266
272
264
266
262
115 115 113 113 Rg Mt
Bh Bh (missed) Bh Bh Bh Bh
Bohrium
7
Chronology of isotope discovery Isotope
Year discovered
discovery reaction
260
2007
209
261
1989
209
262
1981
209
263
unknown
264
1994
209
265
2004
243
266
2004
209
267
2000
249
268
unknown
269
unknown
270
2006
271
unknown
272
2003
Bh Bh Bhg,m Bh Bh Bh Bh Bh Bh Bh Bh Bh Bh
Bi(52Cr,n) Bi(54Cr,2n) Bh(54Cr,n)
Bi(64Ni,n) Am(26Mg,4n) Bi(70Zn,n) Bk(22Ne,5n)
237
[16]
243
[17]
Np(48Ca,3n)
Am(48Ca,3n)
Isomerism in bohrium nuclides 262
Bh
The only confirmed example of isomerism in bohrium is for the isotope 262Bh. Direct production populates two states, a ground state and an isomeric state. The ground state is confirmed as decaying by alpha emission with alpha lines at 10.08,9.82 and 9.76 MeV with a revised half life of 84 ms. The excited state decays by alpha emission with lines at 10.37 and 10.24 MeV with a revised half-life of 9.6 ms.
Chemical yields of isotopes Cold Fusion The table below provides cross-sections and excitation energies for cold fusion reactions producing bohrium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
CN
1n
2n
55
208
263
590 pb , 14.1 MeV
~35 pb
54
209
263
510 pb , 15.8 MeV
~50 pb
52
209
261
59 pb , 15.0 MeV
Mn Cr Cr
Pb Bi Bi
Bh Bh Bh
3n
Bohrium
8
Hot Fusion The table below provides cross-sections and excitation energies for hot fusion reactions producing bohrium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
CN
26
243
271
22
249
271
Mg Ne
Am Bk
Bh
3n +
Bh
4n
5n
+
+
~96 pb
+
External links • • • •
WebElements.com - Bohrium [18] Apsidium - Bohrium [19] Los Alamos National Laboratory - Bohrium Properties of BhO3Cl [21]
[20]
References [1] "GAS CHEMICAL INVESTIGATION OF BOHRIUM (Bh, ELEMENT 107)" (http:/ / www. gsi. de/ informationen/ wti/ library/ scientificreport2000/ Chemistry/ 9/ r_eichler_jb2000. pdf), Eichler et al.., GSI Annual Report 2000. Retrieved on 2008-02-29 [2] [3] [4] [5] [6]
http:/ / www. iupac. org/ publications/ pac/ 1994/ pdf/ 6612x2419. pdf (IUPAC 1994 recomm) http:/ / www. iupac. org/ publications/ pac/ 1997/ pdf/ 6912x2471. pdf (IUPAC 1997 recomm) http:/ / www. nrc-cnrc. gc. ca/ eng/ education/ elements/ el/ bh. html http:/ / www. iupac. org/ publications/ pac/ 1993/ pdf/ 6508x1757. pdf (TWG report) "Identification of element 107 by α correlation chains" (http:/ / www. springerlink. com/ content/ tlx33361417u3n26/ ), Munzenberg et al., Z. Phys. A., 1981, 300, 1. Retrieved on 2008-02-29 [7] "Element 107" (http:/ / www. springerlink. com/ content/ xw1511792061371x/ ), Munzenberg et al., Z. Phys. A., 1989, 333, 2. Retrieved on 2008-02-29 [8] "Entrance Channel Effects in the Production of 262,261Bh" (http:/ / rnc. lbl. gov/ nsd/ annualreport2005/ contributions/ Nelson_LE. pdf), Nelson et al., LBNL repositories 2005. Retrieved on 2008-03-04 [9] "Lightest Isotope of Bh Produced Via the 209Bi(52Cr,n)260Bh Reaction" (http:/ / repositories. cdlib. org/ cgi/ viewcontent. cgi?article=5987& context=lbnl), Nelson et al., LBNL repositories, May 7, 2007. Retrieved on 2008-02-29 [10] "Excitation function for the production of 262Bh (Z=107) in the odd-Z-projectile reaction 208Pb(55Mn, n)" (http:/ / scitation. aip. org/ getabs/ servlet/ GetabsServlet?prog=normal& id=PRVCAN000073000001014611000001& idtype=cvips& gifs=yes), Folden et al., Phys. Rev. C, 73, 014611 (2006). Retrieved on 2008-02-29 [11] "Excitation function for the production of 262Bh (Z=107) in the odd-Z-projectile reaction 208Pb(55Mn, n)" (http:/ / repositories. cdlib. org/ cgi/ viewcontent. cgi?article=3948& context=lbnl), Folden et al., LBNL repositories, May 19, 2005. Retrieved on 2008-02-29 [12] http:/ / www-wnt. gsi. de/ tasca06/ images/ contributions/ TASCA_06_Gates. pdf [13] "New isotope 265Bh" (http:/ / www. ingentaconnect. com/ content/ klu/ 10050/ 2004/ 00000020/ 00000003/ art00006), Gan et al., Eur. Phys. J. A., 2004, 20, 3, 385-387. Retrieved on 2008-02-29 [14] http:/ / xxx. lanl. gov/ PS_cache/ arxiv/ pdf/ 0904/ 0904. 1093v1. pdf [15] "Evidence for New Isotopes of Element 107: 266Bh and 267Bh" (http:/ / prola. aps. org/ abstract/ PRL/ v85/ i13/ p2697_1), Wilk et al., Phys. Rev. Lett., 85, 2697-2700 (2000).Retrieved on 2008-02-29 [16] [17] [18] [19] [20] [21]
see ununtrium see ununpentium http:/ / www. webelements. com/ webelements/ elements/ text/ Bh/ index. html http:/ / www. apsidium. com/ elements/ 107. htm http:/ / periodic. lanl. gov/ elements/ 107. html http:/ / bh-bohrium. info/ properties. html
Article Sources and Contributors
Article Sources and Contributors Bohrium Source: http://en.wikipedia.org/w/index.php?oldid=308845609 Contributors: Ahoerstemeier, AlimanRuna, Andres, Atomica226, BlueEarth, Bryan Derksen, Bsimmons666, Bunny Angel13, CardinalDan, Carnildo, Conversion script, Curps, DO'Neil, Darksun, Darrien, Deanos, Dirac66, Doctorfluffy, Drjezza, Edgar181, Eliyak, Emperorbma, Enceladus, Eras-mus, Femto, Furrykef, Gamma, Goods21, Greatpatton, Guaka, Hede2000, Icairns, Ideyal, J.delanoy, JForget, Joanjoc, John, Jossi, Kalamkaar, Karelj, Kelovy, Kgf0, Kingdon, Kurator, Kurykh, Kwamikagami, Lethalgeek, LiDaobing, Magicindark, Materialscientist, Mav, Merovingian, Muke, N-true, Neitherday, Nergaal, NewEnglandYankee, PierreAbbat, PlatinumX, Poolkris, Razorflame, Rbraunwa, Remember, Reza kalani, Rich Farmbrough, Rifleman 82, Roentgenium111, Romanm, Romanskolduns, Rursus, Samir, Saperaud, Sceptre, Schneelocke, Siffler, Simon J Kissane, Sl, Slashme, Stifynsemons, Tagishsimon, Tarquin, Tetracube, VASANTH S.N., Warut, Yekrats, Yurik, 55 anonymous edits
Image Sources, Licenses and Contributors image:Bh-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Bh-TableImage.png License: GNU Free Documentation License Contributors: Ahoerstemeier, Kwamikagami File:Speaker Icon.svg Source: http://en.wikipedia.org/w/index.php?title=File:Speaker_Icon.svg License: Public Domain Contributors: Blast, G.Hagedorn, Mobius, 2 anonymous edits
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
9
Hassium
1
Hassium bohrium ← hassium → meitnerium Os ↑ Hs ↓ (Upo) Periodic Table Extended Periodic Table General Name, symbol, number Element category Group, Period, Block Appearance
hassium, Hs, 108 transition metals 8, 7, d unknown, probably silvery white or metallic gray
Standard atomic weight Electron configuration Electrons per shell
−1
[277] g·mol 14
[Rn] 5f
6
2
6d 7s
2, 8, 18, 32, 32, 14, 2 Physical properties
Phase
presumably a solid Atomic properties Crystal structure
Oxidation states CAS registry number
unknown 8 54037-57-9 Most-stable isotopes
Hassium
2
Main article: Isotopes of hassium iso
NA
half-life
DM
277
syn
16.5 min
SF
275
syn
0.15 s
α
271
syn
40# s
270
syn
269
Hs ?
DE (MeV)
DP
9.30
271
α
9.27,9.13
267
7.9 s
α
9.02,8.88
266
syn
9.7 s
α
9.21,9.10,8.97
265
267m
syn
0.8 s
α
9.83
263
267
syn
52 ms
α
9.87
263
266
syn
2.3 ms
α
10.18
262
265m
syn
0.75 ms
α
261
265
syn
2.0 ms
α
261
264
syn
~0.8 ms
.5 α
Hs Hs Hs Hs Hs
Hs Hs Hs
Hs Hs
Sg Sg Sg Sg Sg Sg Sg Sg Sg
10.43
260
10.89,10.72,10.57
259
Sg
.5 SF 263
Hs
syn
0.74 ms
α
Sg
References
Hassium (pronounced /ˈhæsi.əm/ ( listen) or English pronunciation: /ˈhɑːsi.əm/[1] ) is a synthetic element in the periodic table that has the symbol Hs and atomic number 108. Its most stable isotope is 277Hs, with a half-life of 16.5 minutes.
Official discovery Hassium was first synthesized in 1984 by a German research team led by Peter Armbruster and Gottfried Münzenberg at the Institute for Heavy Ion Research (Gesellschaft für Schwerionenforschung) in Darmstadt. The team bombarded a lead target with 58Fe nuclei to produce 3 atoms of 265Hs in the reaction: 20882Pb
+
5826Fe
→
265108Hs
+n
The IUPAC/IUPAP Transfermium Working Group (TWG) recognised the GSI collaboration as official discoverers in their 1992 report. [2]
Hassium
3
Naming Element 108 has historically been known as eka-osmium. During the period of controversy over the names of the elements (see element naming controversy) IUPAC adopted unniloctium (symbol Uno) as a temporary element name for this element. The name hassium was proposed by the officially recognised German discoverers in 1992, derived from the Latin name for the German state of Hesse where the institute is located (L. hassia German Hessen). In 1994 a committee of IUPAC recommended that element 108 be named hahnium (Hn)[3] , in spite of the long-standing convention to give the discoverer the right to suggest a name. After protests from the German discoverers, the name hassium (Hs) was adopted internationally in 1997.[4]
Eka-osmium Eka-osmium was a temporary name used to refer to the element that goes under osmium in the periodic table. The name "eka" was used in the same way as in Mendeleev's predicted elements. During the first half of the 20th century, eka-osmium referred to plutonium, because the actinide concept, which postulates the actinides form an inner transition series similar to the lanthanides, had not been proposed yet. Once the actinide concept became widely accepted, eka-osmium started to refer to element 108, now called Hassium, which was discovered in 1984.
Electronic structure Hassium has 6 full shells, 7s+5p+3d+2f=17 full subshells, and 108 orbitals: Bohr model Quantum mechanical model
2, 8, 18, 32, 32, 14, 2 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f146d6
Extrapolated chemical properties of Hassium Oxidation states Element 108 is projected to be the fifth member of the 6d series of transition metals and the heaviest member of group VIII in the Periodic Table, below iron, ruthenium and osmium. The latter two members of the group readily portray their group oxidation state of +8 and this state becomes more stable as the group is descended. Thus hassium is expected to form a stable +8 state. Osmium also shows stable +5, +4 and +3 states with the +4 state the most stable. For ruthenium, the +6, +5 and +3 states are stable with the +3 state being the most stable. Hassium is therefore expected to also show other stable lower oxidation states.
Hassium
4
Chemistry The group VIII elements show a very distinctive oxide chemistry which allows facile extrapolations to be made for hassium. All the lighter members have known or hypothetical tetroxides, MO4. The oxidising power decreases as one descends the group such that FeO4[5] is not known due to an extraordinary electron affinity which results in the formation of the well-known oxo-ion ferrate(VI), FeO42−. Ruthenium tetroxide, RuO4, formed by oxidation of ruthenium(VI) in acid, readily undergoes reduction to ruthenate(VI), RuO42−. Oxidation of ruthenium metal in air forms the dioxide, RuO2. In contrast, osmium burns to form the stable tetroxide, OsO4, which complexes with hydroxide ion to form an osmium(VIII) -ate complex, [OsO4(OH)2]2−. Therefore, eka-osmium properties for hassium should be demonstrated by the formation of a volatile tetroxide HsO4, which undergoes complexation with hydroxide to form a hassate(VIII), [HsO4(OH)2]2−.
Experimental chemistry Gas phase chemistry Hassium is expected to have the electron configuration [Rn]5f14 6d6 7s2 and thus behave as the heavier homolog of osmium (Os). As such, it should form a volatile tetroxide, HsO4, due to the tetrahedral shape of the molecule. The first chemistry experiments were performed using gas thermochromatography in 2001, using 172Os as a reference. During the experiment, 5 hassium atoms were detected using the reaction 248Cm(26Mg,5n)269Hs. The resulting atoms were thermalized and oxidized in a He/O2 mixture to form the oxide. 269108Hs
+ 2 O2 →
269108HsO
4
The measured deposition temperature indicated that hassium(VIII) oxide is less volatile than osmium tetroxide, OsO4, and places hassium firmly in group 8.[6] [7] In order to further probe the chemistry of hassium, scientists decided to assess the reaction between hassium tetroxide and sodium hydroxide to form sodium hassate(VIII), a reaction well-known with osmium. In 2004, scientists announced that they had succeeded in carrying out the first acid-base reaction with a hassium compound:[8] HsO4 + 2 NaOH → Na2[HsO4(OH)2]
Summary of compounds and complex ions Formula
Names(s)
HsO4
hassium tetroxide; hassium(VIII) oxide
Na2[HsO4(OH)2]
sodium hassate(VIII); disodium dihydroxytetraoxohassate(VIII)
History of synthesis of isotopes by cold fusion This section deals with the synthesis of nuclei of hassium by so-called "cold" fusion reactions. These are processes which create compound nuclei at low excitation energy (~10-20 MeV, hence "cold"), leading to a higher probability of survival from fission. The excited nucleus then decays to the ground state via the emission of one or two neutrons only.
Hassium 136
Xe(136Xe,xn)272−xHs
Important future experiments will involve the attempted synthesis of hassium isotopes in this symmetric reaction using the fission fragments. This reaction was carried out at Dubna in 2007 but no atoms were detected, leading to a cross section limit of 1 pb.[9] If confirmed, this would indicate that such symmetric fusion reactions should be modelled as 'hot fusion' reactions rather than 'cold fusion' ones, as first suggested. This would indicate that such reactions will unfortunately have limited use in the synthesis of superheavy elements. 198
Pt(70Zn,xn)268−xHs
This reaction was performed in May 2002 at the GSI. Unfortunately, the experiment was cut short due to a failure of the zinc-70 beam. 208
Pb(58Fe,xn)266−xHs (x=1,2)
This reaction was first reported in 1978 by the team at Dubna. In a later experiment in 1984, using the rotating drum technique, they were able to detect a spontaneous fission activity assigned to 260Sg, daughter of 264Hs. [10] In a repeat experiment in the same year, they applied the method of chemical identification of a descendant to provide support to the synthesis of element 108. They were able to detect several alpha decays of 253Es and 253 Fm, descendants of 265108. In the official discovery of the element in 1984, the team at GSI studied the reaction using the alpha decay genetic correlation method. They were able to positively identify 3 atoms of 265 Hs. [11] After an upgrade of their facilities in 1993, the team repeated the experiment in 1994 and detected 75 atoms of 265Hs and 2 atoms of 264Hs, during the measurement of a partial excitation function for the 1n neutron evaporation channel.[12] The maximum of the 1n channel was measured as 69 pb in a further run in late 1997 in which a further 20 atoms were detected.[13] The discovery experiment was successfully repeated in 2002 at RIKEN (10 atoms) and in 2003 at GANIL (7 atoms). The team at RIKEN further studied the reaction in 2008 in order to conduct first spectroscopic studies of the even-even nucleus 264Hs. 207
Pb(58Fe,xn)265−xHs (x=1)
The use of a Pb-207 target was first used in 1984 at Dubna. They were able to detect the same SF activity as observed in the Pb-208 run and once again assigned it to 260Sg, daughter of 264Hs.[2] The team at GSI first studied the reaction in 1986 using the method of correlation of genetic alpha decays and identified a single atom of 264Hs with a cross section of 3.2 pb.[14] The reaction was repeated in 1994 and the team were able to measure both alpha decay and spontaneous fission for 264Hs. This reaction was studied in 2008 at RIKEN in order to conduct first spectrscopic studies of the even-even nucleus 264Hs.
5
Hassium 208
Pb(56Fe,xn)264−xHs (x=1)
This reaction was studied for the first time in 2008 by the team at LBNL. They were able to produce and identify 6 atoms of the new isotope 263Hs.[15] A few months later, the RIKEN team also published their results on the same reaction.[16] 206
Pb(58Fe,xn)264−xHs (x=1)
This reaction was studied for the first time in 2008 by the team at RIKEN. They were able to identify the new isotope 263Hs.[17] 209
Bi(55Mn,xn)264−xHs
First attempts to synthesise nuclei of element 108 were performed using this reaction by the team at Dubna in 1983. Using the rotating drum technique, they were able to detect a spontaneous fission activity assigned to 255Rf, descendant of the 263108 decay chain. Identical results were measured in a repeat run in 1984.[2] In a subsequent experiment in 1983, they applied the method of chemical identification of a descendant to provide support to the synthesis of element 108. They were able to detect alpha decays from fermium isotopes, assigned as descendants of the decay of 262108. This reaction has not been tried since and 262Hs is currently unconfirmed.[2]
History of synthesis of isotopes by hot fusion This section deals with the synthesis of nuclei of hassium by so-called "hot" fusion reactions. These are processes which create compound nuclei at high excitation energy (~40-50 MeV, hence "hot"), leading to a reduced probability of survival from fission and quasi-fission. The excited nucleus then decays to the ground state via the emission of 3-5 neutrons. 226
Ra(48Ca,xn)274−xHs (x=4)
This reaction was reportedly first studied in 1978 by the team at the Flerov Laboratory of Nuclear Reactions (FLNR) under the leadership of Yuri Oganessian. However, results are not available in the literature.[2] The reaction was repeated at the FLNR in June 2008 and results show that the 4 atoms of the isotope 270Hs were detected with a yield of 9 pb. The decay data for the recently discovered isotope was confirmed, although the alpha energy was slightly higher.[18] In Jan 2009, the team repeated the experiment and a further 2 atoms of 270Hs were detected.[19] 232
Th(40Ar,xn)272−xHs
This reaction was first studied at Dubna in 1987. Detection was by spontaneous fission and no activities were found leading to a calculated cross section limit of 2 pb.[2] 238
U(36S,xn)274−xHs (x=4)
This reaction with the rare and expensive 36S isotope was conducted at the GSI in April-May 2008. Preliminary results show that a single atom of 270Hs was detected with a yield of 0.8 pb. The data confirms the decay properties of 270Hs and 266Sg.[20]
6
Hassium
7
238
U(34S,xn)272−xHs (x=5)
In March 1994, the team at Dubna led by the late Yuri Lazerev announced the detection of 3 atoms of 267Hs from the 5n neutron evaporation channel. [21] The decay properties was confirmed by the team at GSI in their simultaneous study of element 110. 248
Cm(26Mg,xn)274−xHs (x=3,4,5)
Most recently, a GSI-PSI collaboration has studied the nuclear reaction of curium-248 with magnesium-26 ions. Between May 2001 and August 2005, the team has studied the excitation function of the 3n, 4n, and 5n evaporation channels leading to 269Hs, 270Hs, and 271 Hs.[22] [23] The synthesis of the important isotope 270Hs was published in December 2006 by the team of scientists from the Technical University of Munich.[24] It was reported that this isotope decayed by emission of an alpha-particle with an energy of 8.83 MeV and a + + 266 projected half-life of ~22 s, assuming a 0 to 0 ground state decay to Sg using the Viola-Seaborg equation. 248
Cm(25Mg,xn)273−xHs
This new reaction was studied at the GSI in July-August 2006 in a search for the new isotope 268Hs. They were unable to detect any atoms from neutron evaporation and calculated a cross section limit of 1 pb. 249
Cf(22Ne,xn)271−xHs
The team at Dubna studied this reaction in 1983 using detection by spontaneous fission (SF). Several short SF activities were found indicating the formation of nuclei of element 108. [2]
Chemical yields of isotopes The tables below provides cross-sections and excitation energies for reactions producing hassium isotopes directly. Data in bold represent maxima derived from excitation function measurements. + represents an observed exit channel.
Cold fusion Projectile
Target
CN
1n
58
208
266
69 pb, 13.9 MeV
58
207
265
3.2 pb
Fe
Pb
Fe
Pb
Hs Hs
2n 4.5 pb
Hot fusion Projectile
Target
CN
3n
4n
48
226
274
9.0 pb
36
238
274
0.8 pb
34
238
272
Ca S S
Ra U U
Hs Hs Hs
5n
2.5 pb, 50.0 MeV
3n
Hassium
8
26
248
Mg
Cm
274
Hs
2.5 pb
3.0 pb
7.0 pb
Isomerism in hassium isotopes 269
Hs
The direct synthesis of 269Hs has resulted in three alpha lines at 9.21, 9.10, and 8.94 MeV. In the decay of 277112, only 9.21 MeV 269Hs alpha decays have been observed indicating that this decay occurs from an isomeric level. Further research is required to confirm this. 267
Hs
The decay of 267Hs is known to occur by alpha decay with three alpha lines at 9.88, 9.83, and 9.75 MeV and a half-life of 52 ms. In the recent syntheses of 271m,gDs additional activities have been observed. A .94ms activity decaying by 9.83 MeV alpha emission has been observed in addition to longer lived ~.8 s and ~6.0 s activities. Each of these is currently not assigned and confirmed and further research is required to positively identify them. 265
Hs
The synthesis of 265Hs has also provided evidence for two levels. The ground state decays by 10.30 MeV alpha emission with a half-life of 2.0 ms. The isomeric state is placed at 300 keV above the ground state and decays by 10.57 MeV alpha emission with a half-life of .75 ms.
Chronology of isotope discovery Isotope
Year discovered
Discovery reaction
263
2008
208
264
1986
207
265
1984
208
266
2000
207
267
1995
238
268
unknown
269
1996
208
270
2004
248
271
2004
248
272
unknown
273
unknown
274
unknown
275
2003
Hs Hs Hs Hs Hs Hs Hs Hs Hs Hs Hs Hs Hs
Pb(56Fe,n) Pb(58Fe,n) Pb(58Fe,n) Pb(64Ni,n)
[25]
U(34S,5n)
Pb(70Zn,n)
[26]
Cm(26Mg,4n) Cm(26Mg,3n)
242
Pu(48Ca,3n)
[27]
Hassium
9
276
unknown
277
1999?
Hs Hs
244
Pu(48Ca,3n)
[27]
270
Hs: prospects for a deformed doubly-magic nucleus
According to macroscopic-microscopic (MM) theory, Z=108 is a deformed proton magic number. This means that such nuclei are permanently deformed in their ground state but have high, narrow fission barriers to further deformation and hence relatively-long SF partial half-lives. The SF half-lives in this region are typically reduced by a factor of 109 in comparison with those in the vicinity of the spherical doubly-magic nucleus 298114, caused by an increase in the probability of barrier penetration by quantum tunnelling, due to the narrower fission barrier. In addition, N=162 has been calculated as a deformed neutron magic number and hence the nucleus 270Hs has promise as a deformed doubly-magic nucleus. Experimental data from the decay of Z=110 isotopes 271Ds and 273Ds, provides strong evidence for the magic nature of the N=162 sub-shell. The recent synthesis of 269Hs, 270 Hs, and 271Hs also fully support the assignment of N=162 as a magic closed shell.
Evidence for the Z=108 deformed proton shell Evidence for the effect of the Z=108 closed shell in the vicinity of the N=162 shell is limited at this moment in time. This is caused by the low production yields of the isotopes in question and thus poor statistics regarding SF partial half-lives resulting from branching of the decay mode. In the case of the isotonic pair 264Hs and 262Sg (N=156 isotones), the lifetimes and decay modes do not support the stabilizing effect of Z=108 but this is most likely due to a retreat from the N=162 shell. More conclusive evidence would come from the measurement of SF partial half-lives for 266Hs (vs. 264Sg), 268Hs (vs. 266Sg), and especially 270Hs itself (vs 268Sg and 266Rf), although 268Sg and 268Hs are currently unknown and 266Rf has not been produced via alpha decay (which would provide TSF for this N=162 isotone). Analysis of partial SF half-lives of nuclei with Z>108 (e.g. 272Ds) would also help to confirm the Z=108 closed shell. It should be noted that whilst 270Hs is expected to be a doubly-magic nucleus, it is not expected to have the longest half-life in this region of the periodic table. The reason is that whilst the N=162 shell staves off fission, alpha decay will predominate. As an example, the nucleus 268Sg (Z=106,N=162) is calculated to have a halflife of the order of two hours. However, recent data from the decay of 264Sg (TSF = 70 ms) and 266Sg (TSF = 360 ms) indicate that the influence of the N=162 shell for seaborgium isotopes against fission is some 1–2 orders of magnitude overestimated, so 268Sg may in fact decay by SF will a short half life of ~5 s. The recently-synthesized nucleus 268Db (TSF = 29 h) has such a long half-life because the presence of both the odd proton and odd neutron hinder SF, relative to neighbouring even-even nuclei.
Hassium
10
Unconfirmed isotopes 277
Hs
An isotope assigned to 277Hs has been observed on two occasions decaying by SF with a long half-life of ~12 minutes. The isotope is not observed in the decay of ground state 281Ds but is observed in the decay from an unconfirmed isomeric level. The half-life is very long for the ground state and it is possible that it belongs to an isomeric level in 277Hs. Further research is required to confirm this result.
Retracted isotopes 273
Hs
The claimed synthesis of element 118 by LBNL in 1999 involved the intermediate 273Hs. This isotope was claimed to decay by 9.78 and 9.47 MeV alpha emission with a half-life of 1.2 s. The claim to discovery of 293118 was retracted and this hassium isotope is currently unknown.
Theoretical calculations on evaporation cross sections The below table contains various targets-projectile combinations for which calculations have provided estimates for cross section yields from various neutron evaporation channels. The channel with the highest expected yield is given. DNS = Di-nuclear system ; σ = cross section Target 136
Xe
238
U
Projectile 136
Xe
34
S
CN
Channel (product)
~σ
max
Model
Ref
272
1-4n (271-268Hs)
10-6 pb
DNS
[28]
272
4n (268Hs)
10 pb
DNS
[28]
Hs Hs
Current experiments The team at GSI studied the synthesis of hassium isotopes by hot fusion in Jan-Feb 2009, using the reaction 238U(34S,xn).[29] Results are not yet available.
Future experiments Spectroscopy Scientists at the GSI are planning to search for K-isomers in 270Hs using the reaction 226 Ra(48Ca,4n) in 2010. They will use the new TASISpec method developed alongside the introduction of the new TASCA facility at the GSI.[30] In addition, they also hope to study the spectroscopy of 269Hs, 265Sg and 261Rf, using the reaction 248Cm(26Mg,5n) or 226Ra(48Ca,5n). This will allow them to determine the level structure in 265Sg and 261Rf and attempt to give spin and parity assignments to the various proposed isomers.[31]
Hassium
11
Chemistry The team from the universität Mainz are planning to study the electrodeposition of hassium atoms using TASCA at the GSI. The current aim is to use the reaction 226 Ra(48Ca,4n)270Hs.[32] In addition, scientists at the GSI are hoping to utilize the new TASCA facility to study the synthesis and properties of the hassium(II) compound, hassocene, Hs(cp)2 using the reaction 226Ra(48Ca,xn).[33]
External links • WebElements.com: Hassium • Apsidium: Hassium 108 [35]
[34]
References [1] hassium at Dictionary.com (http:/ / dictionary. reference. com/ browse/ hassium) [2] "Discovery of the transfermium elements" (http:/ / iupac. org/ publications/ pac/ 1993/ pdf/ 6508x1757. pdf), IUPAC Technical report, Pure & Appl. Chem., Vol. 65, No. 8, pp. 1757-1814,1993. Retrieved on 2008-03-07 [3] http:/ / iupac. org/ publications/ pac/ 1994/ pdf/ 6612x2419. pdf (IUPAC 1994 recomm) [4] http:/ / iupac. org/ publications/ pac/ 1997/ pdf/ 6912x2471. pdf (IUPAC 1997 recomm) [5] "FeO4: A unique example of a closed-shell cluster mimicking a superhalogen" (http:/ / prola. aps. org/ abstract/ PRA/ v59/ i5/ p3681_1), Gutsev et al., Phys. Rev. A , 59, 3681- 3684 (1999). Retrieved on 2008-03-01 [6] http:/ / lch. web. psi. ch/ pdf/ anrep01/ B-03heavies. pdf (Investigation of Hassium) [7] " "Chemistry of Hassium" (http:/ / www. gsi. de/ documents/ DOC-2003-Jun-29-2. pdf)" (PDF). Gesellschaft für Schwerionenforschung mbH. 2002. . Retrieved 2007-01-31. [8] http:/ / www. gsi. de/ informationen/ wti/ library/ scientificreport2003/ files/ 172. pdf (CALLISTO result) [9] http:/ / www1. jinr. ru/ Reports/ 2007/ english/ 06_flnr_e. pdf [10] "On the stability of the nuclei of element 108 with A=263–265" (http:/ / springerlink. com/ content/ p1107209h687r205/ ), Oganessian et al.., Z. Phys. A., 1984, 319, 2. Retrieved on 2008-03-01 [11] "The identification of element 108" (http:/ / springerlink. com/ content/ n752260210006404/ ), Munzenberg et al., Z. Phys. A., 1984, 317, 2. Retrieved on 2008-03-01 [12] "New elements-approaching Z=114" (http:/ / www. iop. org/ EJ/ abstract/ 0034-4885/ 61/ 6/ 002), S.Hoffmann, Rep. Prog. Phys.,1998, 61, 639-689.Retrieved on 2008-03-01 [13] "Excitation function for the production of 265108 and 266109" (http:/ / springerlink. com/ content/ b0lj7nyrvh6ugujy/ ), Hofmann et al.., Z. Phys. A., 1997, 358, 4. Retrieved on 2008-03-01 [14] "Evidence for 264108, the heaviest known even-even isotope" (http:/ / springerlink. com/ content/ rn11r8qh66660jnp/ ), Munzenberg et al., Z. Phys. A., 1986, 324, 4. Retrieved on 2008-03-01 [15] Dragojevic et al., "New Isotope 263108" (http:/ / link. aps. org/ doi/ 10. 1103/ PhysRevC. 79. 011602), Phys. Rev. C 2009, 79, 011602, 2008-03-24. [16] Kaji et al., "Production and Decay Properties of 263108" (http:/ / jpsj. ipap. jp/ link?JPSJ/ 78/ 035003), J. Phys. Soc. Jpn. 2009, 78, 035003, 2008-03-24. [17] http:/ / 159. 93. 28. 88/ linkc/ flnr_presentations/ Mendeleev%20simposium/ Morita. ppt [18] http:/ / www1. jinr. ru/ Reports/ 2008/ english/ 06_flnr_e. pdf [19] http:/ / www. np. ph. bham. ac. uk/ iop09/ iop_talks/ 07_04_09_Parallel_2/ Y%20Tsyganov. ppt [20] http:/ / www. gsi. de/ informationen/ wti/ library/ scientificreport2008/ PAPERS/ NUSTAR-SHE-02. pdf [21] "New Nuclide 267108 Produced by the 238U + 34S Reaction" (http:/ / prola. aps. org/ abstract/ PRL/ v75/ i10/ p1903_1), Lazarev et al., Phys. Rev. Lett., 75, 1903-1906 (1995). Retrieved on 2008-03-01 [22] "Decay properties of 269Hs and evidence for the new nuclide 270Hs" (http:/ / lch. web. psi. ch/ pdf/ anrep01/ B-01heavies. pdf), Turler et al., GSI Annual Report 2001. Retrieved on 2008-03-01 [23] http:/ / www2. ha. physik. uni-muenchen. de/ heaviest_atoms/ talks/ Dvorak. pdf (269-271Hs) [24] "Doubly magic 270Hs" (http:/ / www. gsi. de/ informationen/ wti/ library/ scientificreport2006/ PAPERS/ NUSTAR-SHE-06. pdf), Turler et al.., GSI report, 2006. Retrieved on 2008-03-01 [25] see darmstadtium [26] see ununbium [27] see ununquadium [28] http:/ / arxiv. org/ pdf/ 0904. 2994
Hassium [29] [30] [31] [32] [33] [34] [35]
http:/ / www. gsi. de/ documents/ DOC-2007-Mar-178-1. pdf (238U + 34S proposal) http:/ / www-win. gsi. de/ tasca08/ contributions/ TASCA08_Cont_Andersson. pdf http:/ / www-win. gsi. de/ tasca08/ contributions/ TASCA08_Cont_Yakushev2. pdf http:/ / www-win. gsi. de/ tasca08/ contributions/ TASCA08_Cont_Kratz. pdf http:/ / www-win. gsi. de/ tasca08/ contributions/ TASCA08_Cont_Duellmann1. pdf http:/ / www. webelements. com/ webelements/ elements/ text/ Hs/ index. html http:/ / www. apsidium. com/ elements/ 108. htm
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Article Sources and Contributors
Article Sources and Contributors Hassium Source: http://en.wikipedia.org/w/index.php?oldid=308003681 Contributors: Ahoerstemeier, Alechemista, Alex43223, AlimanRuna, Allstarecho, Antandrus, Benbest, BlueEarth, Bryan Derksen, C'est moi, CUSENZA Mario, Can't sleep, clown will eat me, Carnildo, Conversion script, Corpx, Darktemplar, Darrien, Dead3y3, Deanos, Dharumanyo1, Dimotika, Drjezza, Droog Andrey, EPO, Edgar181, Emperorbma, Eras-mus, Femto, Finlay McWalter, Fonzy, GPHemsley, Gamma, Greatpatton, Gurch, Icairns, Icek, Ideyal, J.delanoy, JWBE, John, Jok2000, Kalamkaar, Karelj, Karl-Henner, Kgf0, Kilo-Lima, Kingdon, Kurator, Kwamikagami, Lars T., LiDaobing, Markus liverpool, Mav, Mets501, Mitchandre, Muke, Nergaal, Nihiltres, Odie5533, Oliverdl, PTSE, Pilotguy, Poolkris, Pras, Remember, Rich Farmbrough, Rifleman 82, Rjwilmsi, Robertb-dc, Roentgenium111, RoryReloaded, Rursus, Saganatsu, Schneelocke, Simon J Kissane, Sl, Splash, Spug, Squids and Chips, Stevey7788, Stifynsemons, Stone, SummerPhD, Tagishsimon, Tavilis, Tetracube, Timc, Titoxd, Tonyrex, VASANTH S.N., Vsmith, Vuo, Yekrats, Zetawoof, Zotel, 118 anonymous edits
Image Sources, Licenses and Contributors image:Hs-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Hs-TableImage.png License: GNU Free Documentation License Contributors: Dodo, Kwamikagami, Paddy, Saperaud File:Speaker Icon.svg Source: http://en.wikipedia.org/w/index.php?title=File:Speaker_Icon.svg License: Public Domain Contributors: Blast, G.Hagedorn, Mobius, 2 anonymous edits
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13
Meitnerium
1
Meitnerium hassium ← meitnerium → darmstadtium Ir ↑ Mt ↓ (Upe) Periodic Table Extended Periodic Table General Name, symbol, number Element category Group, Period, Block Appearance
meitnerium, Mt, 109 transition metals 9, 7, d probably light silver green metallic
Standard atomic weight Electron configuration
Electrons per shell Phase
−1
[276] g·mol
14
2, 8, 18, 32, 32, 15, 2 presumably a solid
CAS registry number
7
2
perhaps [Rn] 5f 6d 7s (guess based on iridium)
54038-01-6 Most-stable isotopes
Meitnerium
2
Main article: Isotopes of meitnerium iso
NA
half-life
279
syn
6 min (est.)
278
syn
30 min (est.)
277
syn
1 min (est.)
276
syn
0.72 s
275
syn
9.7 ms
274
syn
270m
Mt ?
Mt
Mt Mt Mt
DM
DP
9.71
272
α
10.33
271
0.44 s
α
9.76
270
syn
1.1 s
α
270g
syn
5 ms
α
10.03
266
268
syn
42 ms
α
10.26,10.10
264
266
syn
1.7 ms
α
11.00
262
Mt Mt Mt
Mt Mt
α
DE (MeV)
Bh Bh Bh
266
Bh Bh Bh Bh
References [1] Meitnerium (pronounced /maɪtˈnɜriəm/ ( listen)) is a chemical element in the periodic table that has the symbol Mt and atomic number 109.
Mt is a synthetic element whose most stable known isotope is Mt-276, with a half-life of a 0.7 s.
Discovery profile Meitnerium was first synthesized on August 29, 1982 by a German research team led by Peter Armbruster and Gottfried Münzenberg at the Institute for Heavy Ion Research (Gesellschaft für Schwerionenforschung) in Darmstadt.[2] The team bombarded a target of bismuth-209 with accelerated nuclei of iron-58 and detected a single atom of the isotope meitnerium-266: 20983Bi
+
5826Fe
→
266109Mt
+n
Naming Historically, element 109 has been referred to as eka-iridium. The name meitnerium (Mt) was suggested in honor of the Austrian physicist Lise Meitner. In 1997, the name was officially adopted by the IUPAC.
Electronic structure Meitnerium is element 109 in the Periodic Table. The two forms of the projected electronic structure are: Bohr model
2, 8, 18, 32, 32, 15, 2
Meitnerium
Quantum mechanical [3] model
3
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f146d7
Extrapolated chemical properties of meitnerium Physical properties Mt should be a very heavy metal with a density around 30 g/cm3 (Co: 8.9, Rh: 12.5, Ir: 22.5) and a high melting point around 2600-2900°C (Co: 1480, Rh: 1966, Ir: 2454). It should be very corrosion resistant more than Ir which is already the most corrosion resistant metal.
Oxidation states Meitnerium is projected to be the sixth member of the 6d series of transition metals and the heaviest member of group 9 in the Periodic Table, below cobalt, rhodium and iridium. This group of transition metals is the first to show lower oxidation states and the +9 state is not known. The latter two members of the group show a maximum oxidation state of +6, whilst the most stable states are +4 and +3 for iridium and +3 for rhodium. Meitnerium is therefore expected to form a stable +3 state but may also portray stable +4 and +6 states.
Chemistry The +VI state in group 9 is known only for the fluorides which are formed by direct reaction. Therefore, meitnerium should form a hexafluoride, MtF6. This fluoride is expected to be more stable than iridium(VI) fluoride, as the +6 state becomes more stable as the group is descended. In combination with oxygen, rhodium forms Rh2O3 whilst iridium is oxidised to the +4 state in IrO2. Meitnerium may therefore show a dioxide, MtO2, if eka-iridium reactivity is shown. The +3 state in group 9 is common in the trihalides (except fluorides) formed by direct reaction with halogens. Meitnerium should therefore form MtCl3, MtBr3 and MtI3 in an analogous manner to iridium.
History of synthesis of isotopes in cold fusion This section deals with the synthesis of nuclei of meitnerium by so-called "cold" fusion reactions. These are processes which create compound nuclei at low excitation energy (~10-20 MeV, hence "cold"), leading to a higher probability of survival from fission. The excited nucleus then decays to the ground state via the emission of one or two neutrons only. 209
Bi(58Fe,xn)267-xMt (x=1)
The first success in this reaction was in 1982 by the GSI team in their discovery experiment with the identification of a single atom of 266Mt in the 1n neutron evaporation channel.[2] The GSI team used the parent-daughter correlation technique. After an initial failure in 1983, in 1985 the team at the FLNR, Dubna, observed alpha decays from the descendant 246 Cf indicating the formation of meitnerium. The GSI synthesised a further 2 atoms of 266 Mt in 1988 and continued in 1997 with the detection of 12 atoms during the measurement of the 1n excitation function. [4] [5]
Meitnerium
4
208
Pb(59Co,xn)267-xMt (x=1)
This reaction was first studied in 1985 by the team in Dubna. They were able to detect the alpha decay of the descendant 246Cf nuclei indicating the formation of meitnerium atoms. In 2007, in a continuation of their study of the effect of odd-Z projectiles on yields of evaporation residues in cold fusion reactions, the team at LBNL synthesised 266Mt and were able to correlate the decay with known daughters.[6] 181
Ta(86Kr,xn)267-xMt
There are indications that this cold fusion reaction using a tantalum target was attempted in August 2001 at the GSI. No details can be found suggesting that no atoms of meitnerium were detected.
History of synthesis by hot fusion reactions 238
U(37Cl,xn)275-xMt
In 2002-2003, the team at LBNL attempted the above reaction in order to search for the isotope 271Mt with hope that it may be sufficiently stable to allow a first study of the chemical properties of meitnerium. Unfortunately, no atoms were detected and a cross section limit of 1.5 pb was measured for the 4n channel at the projectile energy used. [7] 254
Es(22Ne,xn)276-xMt
Attempts to produce long-living isotopes of meitnerium were first performed by Ken Hulet at the Lawrence Livermore National Laboratory (LLNL) in 1988 using the asymmetric hot fusion reaction above. They were unable to detect any product atoms and established a cross section limit of 1 nb.[8]
Synthesis of isotopes as decay products Isotopes of meitnerium have also been detected in the decay of heavier elements. Observations to date are shown in the table below: Evaporation Residue
Observed Mt isotope
288
276
287
275
282
274
278
270
272
268
115 115 113 113 Rg
Mt Mt Mt Mt Mt
Meitnerium
5
Chronology of isotope discovery Isotope
Year discovered
discovery reaction 209
[2]
209
[9]
209
[10]
266
1982
267
unknown
268
1994
269
unknown
270
2004
271
unknown
272
unknown
273
unknown
274
2006
237
275
2003
243
276
2003
243
Mt Mt Mt Mt Mt Mt Mt Mt Mt Mt Mt
Bi(58Fe,n)
Bi(64Ni,n)
Bi(70Zn,n)
[10]
Np(48Ca,3n)
[11] Am(48Ca,4n) [11] Am(48Ca,3n)
Chemical yields of isotopes Cold Fusion The table below provides cross-sections and excitation energies for cold fusion reactions producing meitnerium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
CN
1n
58
209
267
7.5 pb
59
208
267
2.6 pb , 14.9 MeV
Fe Co
Bi Pb
Mt Mt
2n
3n
Isomerism in meitnerium nuclides 270
Mt
Two atoms of 270Mt have been identified in the decay chains of 278113. The two decays have very different lifetimes and decay energies and are also produced from two apparently different isomers in 274Rg. The first isomer decays by emission of an 10.03 MeV alpha particle with a lifetime 7.2 ms. The other decays by emitting an alpha particle with a lifetime of 1.63 s. An assignment to specific levels is not possible with the limited data available. Further research is required.
Meitnerium
6
268
Mt
The alpha decay spectrum for 268Mt appears to be complicated from the results of several experiments. Alpha lines of 10.28,10.22 and 10.10 MeV have been observed. Half-lives of 42 ms, 21 ms and 102 ms have been determined. The long-lived decay is associated with alpha particles of energy 10.10 MeV and must be assigned to an isomeric level. The discrepancy between the other two half-lives has yet to be resolved. An assignment to specific levels is not possible with the data available and further research is required.
Future Experiments The team at RIKEN, Japan, have indicated that as part of their ongoing studies using targets, they may study the new reaction 248Cm(27Al,xn) in the future.
248
Cm
External links • WebElements.com - Meitnerium • Apsidium - Meitnerium [13]
[12]
References [1] Prof S.Hofmann (private communication) [2] "Observation of one correlated α-decay in the reaction 58Fe on 209Bi→267109" (http:/ / www. springerlink. com/ content/ q4p6m31747740541/ ), Gottfried Munzenberg et al., Z. Phys. A., 1982, 309, 1. Retrieved on 2008-03-01 [3] "Dirac-Hartree-Fock studies of X-ray transitions in meitnerium" (http:/ / epja. edpsciences. org/ index. php?option=article& access=standard& Itemid=129& url=/ articles/ epja/ abs/ 2008/ 05/ 10050_2008_Article_100707/ 10050_2008_Article_100707. html), Christian Thierfelder, Peter Schwerdtfeger, Fritz Peter Heßberger and Sigurd Hofmann , Eur. Phys. J. A, 2008, 36 227 Retrieved on 2008-05-16 [4] "New results on element 109" (http:/ / www. springerlink. com/ content/ x54725101x21h263/ ), Gottfried Munzenberg et al., Z. Phys. A., 1988, 330, 4. Retrieved on 2008-03-01 [5] "Excitation function for the production of 265108 and 266109" (http:/ / www. springerlink. com/ content/ b0lj7nyrvh6ugujy/ ), Sigurd Hofmann et al., Z. Phys. A., 1997, 358, 4. Retrieved on 2008-03-01 [6] Nelson et al. (2009). " Comparison of complementary reactions in the production of Mt (http:/ / scitation. aip. org/ getabs/ servlet/ GetabsServlet?prog=normal& id=PRVCAN000079000002027605000001& idtype=cvips& gifs=yes)". Physical Rev. C 79: 027605. . [7] "The search for 271Mt via the reaction 238U + 37Cl" (http:/ / www. gsi. de/ informationen/ wti/ library/ scientificreport2003/ files/ 2. pdf), Zielinski et al.., GSI Annual report, 2003. Retrieved on 2008-03-01 [8] see reference 4 for reference to an internal report from LLNL [9] see roentgenium for details [10] see ununtrium for details [11] see ununpentium for details [12] http:/ / www. webelements. com/ webelements/ elements/ text/ Mt/ index. html [13] http:/ / www. apsidium. com/ elements/ 109. htm
Article Sources and Contributors
Article Sources and Contributors Meitnerium Source: http://en.wikipedia.org/w/index.php?oldid=308534433 Contributors: AKM, Ahoerstemeier, AlimanRuna, Andres, Beetstra, Benbest, Blurpeace, Bryan Derksen, Carnildo, CatherineMunro, Celarnor, Conversion script, Darrien, Drjezza, Drjezzab, Durova, EPO, Edgar181, Emperorbma, Eras-mus, Femto, Fonzy, Greatpatton, Herbythyme, Icairns, Ideyal, John, Karelj, Karl-Henner, Kingdon, Kurator, Kwamikagami, Lightmouse, Mav, Neitherday, Nergaal, Poolkris, Pras, Psiphiorg, Red Director, Remember, Reza kalani, Rich Farmbrough, Rifleman 82, Roentgenium111, Romanskolduns, Saperaud, Schneelocke, Simon J Kissane, Skatebiker, Sl, Smokefoot, Somethingironic, Stifynsemons, Stone, Tagishsimon, Tetracube, Thricecube, Tomaxer, Tomchiukc, VASANTH S.N., Vuo, Warut, Xezbeth, Yekrats, 73 anonymous edits
Image Sources, Licenses and Contributors image:Mt-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Mt-TableImage.png License: GNU Free Documentation License Contributors: Dodo, Kwamikagami, Paddy, Saperaud File:Speaker Icon.svg Source: http://en.wikipedia.org/w/index.php?title=File:Speaker_Icon.svg License: Public Domain Contributors: Blast, G.Hagedorn, Mobius, 2 anonymous edits
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
7
Darmstadtium
1
Darmstadtium meitnerium ← darmstadtium → roentgenium Pt ↑ Ds ↓ (Uhn) Periodic Table Extended Periodic Table General Name, symbol, number Element category Group, Period, Block Appearance
darmstadtium, Ds, 110 transition metals 10, 7, d unknown, probably silvery white or metallic gray
Standard atomic weight Electron configuration
Electrons per shell Phase
−1
[281] g·mol
14
2, 8, 18, 32, 32, 17, 1 presumably a solid
CAS registry number
9
1
perhaps [Rn] 5f 6d 7s (in analogy to platinum)
54083-77-1 Most-stable isotopes
Darmstadtium
2
Main article: Isotopes of darmstadtium iso
NA
half-life
DM
281
syn
11 s
SF
279
syn
0.20 s
10% α
Ds Ds
DE (MeV)
DP
9.70
275
11.14
269
Hs
90% SF syn
170 ms
271m
Ds
syn
69 ms
α
10.71
267
271g
Ds
syn
1.63 ms
α
10.74,10.69
267
270m
Ds
syn
6 ms
α
12.15,11.15,10.95
266
270g
Ds
syn
0.10 ms
α
11.03
266
269
syn
0.17 ms
α
11.11
265
267
syn
0.004 ms
273
Ds
Ds Ds ?
α
Hs Hs Hs Hs Hs Hs
References [1] Darmstadtium (pronounced /dɑrmˈʃtætiəm/ ( listen)), formerly known as ununnilium, is a chemical element with the symbol Ds and atomic number 110. This synthetic element is one of the so-called super-heavy atoms. It decays quickly: Heavier isotopes of darmstadtium have half-lives on the order of ten seconds.
Official discovery Darmstadtium was first created on November 9, 1994 at the Gesellschaft für Schwerionenforschung (GSI) in Arheilgen, a northern suburb of Darmstadt, Germany by Peter Armbruster and Gottfried Münzenberg, under the direction of professor Sigurd Hofmann. Four atoms of it were detected by a nuclear fusion reaction caused by bombarding a lead-208 target with nickel-62 ions: [2] 20882Pb
+
6228Ni
→
269110Ds
+
10n
In the same series of experiments, the same team also carried out the reaction using heavier nickel-64 ions. During two runs, 9 atoms of 271Ds were convincingly detected by correlation with known daughter decay properties: [3] 20882Pb
+
6428Ni
→
271110Ds
+
10n
The IUPAC/IUPAP Joint Working Party (JWP) recognised the GSI team as discoverers in their 2001 report.[4]
Darmstadtium
3
Proposed names Element 110 was first given the temporary name ununnilium (/ˌjuːnəˈnɪliəm/ or /ˌʌnəˈnɪliəm/[5] , symbol Uun). Once recognized as discoverers, the team at GSI considered the names darmstadtium (Ds) and wixhausium (Wi) for element 110. They decided on the former and named the element after the city near the place of its discovery, Darmstadt and not the suburb Wixhausen itself. The new name was officially recommended by IUPAC on August 16, 2003. The element has also earned the nickname of policium, because the telephone number of the police is 110 within Germany.
Electronic structure Darmstadtium is element 110 in the Periodic Table. The two forms of the projected electronic structure are: Bohr model Quantum mechanical model
2, 8, 18, 32, 32, 17, 1 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s15f146d9
Extrapolated chemical properties of eka-platinum/dvi-palladium Oxidation states Element 110 is projected to be the eighth member of the 6d series of transition metals and the heaviest member of group 10 in the Periodic Table, below nickel, palladium and platinum. The highest confirmed oxidation state of +6 is shown by platinum whilst the +4 state is stable for both elements. Both elements also possess a stable +2 state. Darmstadtium is therefore predicted to show oxidation states +6, +4 and +2.
Chemistry High oxidation states are expected to become more stable as the group is descended, so darmstadtium is expected to form a stable hexafluoride, DsF6, in addition to DsF5 and DsF4. Halogenation should result in the formation of the tetrahalides, DsCl4, DsBr4 and DsI4. Like other Group 10 elements, darmstadtium can be expected to have notable hardness and catalytic properties.
Darmstadtium
History of synthesis of isotopes by cold fusion This section deals with the synthesis of nuclei of darmstadtium by so-called "cold" fusion reactions. These are processes which create compound nuclei at low excitation energy (~10-20 MeV, hence "cold"), leading to a higher probability of survival from fission. The excited nucleus then decays to the ground state via the emission of one or two neutrons only. 208
Pb(64Ni,xn)272-xDs (x=1)
This reaction was first studied by scientists at GSI in 1986, without success. A cross section limit of 12 pb was calculated. After an upgrade of their facilities, they successfully detected 9 atoms of 271Ds in two runs in 1994 as part of their discovery experiments on element 110.[3] This reaction was successfully repeated in 2000 by GSI (4 atoms), in 2000 [6] [7] and 2004 [8] [9] by LBNL (9 atoms in total) and in 2002 by RIKEN (14 atoms).[10] The summation of the data allowed a measurement of the 1n neutron evaporation excitation function. 207
Pb(64Ni,xn)271-xDs (x=1)
In addition to the official discovery reactions, in October-November 2000, the team at GSI also studied the reaction using a Pb-207 target in order to search for the new isotope 270Ds. They succeeded in synthesising 8 atoms of 270Ds, relating to a ground state isomer, 270gDs, and a high-spin K-isomer, 270mDs. [11] 208
Pb(62Ni,xn)270-xDs (x=1)
The GSI team studied this reaction in 1994 as part of their discovery experiment. Three atoms of 269Ds were detected.[2] A fourth decay chain was measured but subsequently retracted. 209
Bi(59Co,xn)268-xDs
This reaction was first studied by the team at Dubna in 1986. They were unable to detect any product atoms and measured a cross section limit of 1 pb. In 1995, the team at LBNL reported that they had succeeded in detecting a single atom of 267Ds from the 1n neutron evaporation channel. However, several decays were missed and further research is required to confirm this discovery. [12]
History of synthesis of isotopes by hot fusion This section deals with the synthesis of nuclei of darmstadtium by so-called "hot" fusion reactions. These are processes which create compound nuclei at high excitation energy (~40-50 MeV, hence "hot"), leading to a reduced probability of survival from fission. The excited nucleus then decays to the ground state via the emission of 3-5 neutrons. Fusion reactions utilizing 48Ca nuclei usually produce compound nuclei with intermediate excitation energies (~30-35 MeV) and are sometimes referred to as "warm" fusion reactions. This leads, in part, to relatively high yields from these reactions.
4
Darmstadtium
5
232
Th(48Ca,xn)280-xDs
The synthesis of element 110 by hot fusion pathways was first attempted in 1986 by the team at Dubna. Using the method of detection of spontaneous fission, they were unable to measure any SF activities and calculated a cross section limit of 1 pb for the decay mode. In three separate experiments between November 1997 and October 1998, the same team re-studied this reaction as part of their new 48Ca program on the synthesis of superheavy elements. Several SF activities with relatively long half-lives were detected and tentatively assigned to decay of the daughters 269Sg or 265Rf, with a cross section of 5 pb. These observations have not been confirmed and the results are taken as only an indication for the synthesis of darmstadtium in this reaction. 232
Th(44Ca,xn)276-xDs
This reaction was attempted in 1986 and 1987 by the Dubna team. In both experiments, a 10 ms SF activities was measured and assigned to 272Ds, with a calculated cross section of 10 pb. This activity is currently not thought to be due to a darmstadtium isotope. 238
U(40Ar,xn)278-xDs
This reaction was first attempted by the Dubna team in 1987. Only spontaneous fission from the transfer products 240mfAm and 242mfAm were observed and the team calculated a cross section limit of 1.6 pb. The team at GSI first studied this reaction in 1990. Once again, no atoms of element 110 could be detected. In August 2001, the GSI repeated reaction, without success, and calculated a cross section limit of 1.0 pb. 236
U(40Ar,xn)276-xDs
This reaction was first attempted by the Dubna team in 1987. No spontaneous fission was observed. 235
U(40Ar,xn)275-xDs
This reaction was first attempted by the Dubna team in 1987. No spontaneous fission was observed. It was further studied in 1990 by the GSI team. Once again, no atoms were detected and a cross section limit of 21 pb was calculated. 233
U(40Ar,xn)273-xDs
This reaction was first studied in 1990 by the GSI team. No atoms were detected and a cross section limit of 21 pb was calculated. 244
Pu(34S,xn)278-xDs (x=5)
In September 1994 the team at Dubna detected a single atom of 273Ds, formed in the 5n neutron evaporation channel. The measured cross section was just 400 fb. [13]
Synthesis of isotopes as decay products Isotopes of darmstadtium have also been detected in the decay of heavier elements. Observations to date are shown in the table below: Evaporation Residue
Observed Ds isotope
Darmstadtium
293
289
291
287
116 , 116 ,
6
281
114 114 ,
Ds
283
279
112
Ds
277
273
112
Ds
In some experiments, the decay of 293116 and 289114 produced an isotope of darmstadtium decaying by emission of an 8.77 MeV alpha particle with a half life of 3.7 minutes. Although unconfirmed, it is highly possible that this activity is associated with a meta-stable isomer, namely 281mDs.
Spectroscopy of darmstadtium isotopes 270
Ds
This is the current partial decay level scheme for 270Ds proposed following the work of Hofmann et al. in 2000 at GSI
Chronology of isotope discovery Isotope
Year discovered
discovery reaction
267
1994
268
unknown
269
1994
208
270
2000
207
Ds ?? Ds Ds Dsg,m
209
Bi(59Co,n)
Pb(62Ni,n) Pb(64Ni,n)
Darmstadtium
7
271
1994
272
unknown
273
1996
274
unknown
275
unknown
276
unknown
277
1997
278
unknown
279
2002
280
unknown
Dsg,m Ds Ds Ds Ds Ds Ds ?? Ds Ds Ds
281a
Ds
281b
Ds ?
208
Pb(64Ni,n)
244
Pu(34S,5n)
232
Th(48Ca,3n)
244
[14]
1999
244
[14]
1999
244
[14]
Pu(48Ca,5n)
Pu(48Ca,3n) Pu(48Ca,3n)
Chemical yields of isotopes Cold fusion The table below provides cross-sections and excitation energies for cold fusion reactions producing darmstadtium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
CN
1n
62
208
270
3.5 pb
64
208
272
15 pb , 9.9 MeV
Ni
Pb
Ni
Pb
Ds Ds
2n
3n
Isomerism in darmstadtium nuclides 281
Ds
The production of 281Ds by the decay of 289114 or 293116 has produced two very conflicting decay modes. The most common and readily confirmed mode is SF with a half-life of 11 s. A much rarer and hitherto unconfirmed mode is alpha decay by emission of an 8.77 MeV alpha particle with an observed half-life of ~3.7 m. This decay is associated with a unique decay pathway from the parent nuclides and must be assigned to an isomeric level. The half-life suggests that it must be assigned to an isomeric state but further research is required to confirm these reports.
Darmstadtium 271
Ds
Decay data from the direct synthesis of 271Ds clearly indicates the presence of two alpha groups. The first has alpha lines at 10.74 and 10.69 MeV with a half-life of 1.63 ms. The other has a single alpha line at 10.71 MeV with a half-life of 69 ms. The first has been assigned to the ground state and the latter to an isomeric level. It has been suggested that the closeness of the alpha decay energies indicates that the isomeric level may decay primarily by delayed gamma emission to the ground state, resulting in an identical measured alpha energy and a combined half-life for the two processes. 270
Ds
The direct production of 270Ds has clearly identified two alpha groups belonging to two isomeric levels. The ground state decays into the ground state of 266Hs by emitting an 11.03 MeV alpha particle with a half-life of 0.10 ms. The isomeric level decays by alpha emission with alpha lines at 12.15,11.15 and 10.95 MeV with a half-life of 6 ms. The 12.15 MeV has been assigned as decay into the ground state of 266Hs indicating that this high spin K-isomer lies at 1.12 MeV above the ground state.
Retracted isotopes 280
Ds
The first synthesis of element 114 resulted in two atoms assigned to 288114, decaying to the 280 Ds which underwent spontaneous fission. The assignment was later changed to 289114 and the darmstadtium isotope to 281Ds. Hence, 280Ds is currently unknown. 277
Ds
In the claimed synthesis of 293118 in 1999, the isotope 277Ds was identified as decaying by 10.18 MeV alpha emission with a half-life of 3.0 ms. This claim was retracted in 2001 and thus this darmstadtium isotope is currently unknown or unconfirmed.[15] 273m
Ds
In the synthesis of 277112 in 1996 by GSI (see ununbium), one decay chain proceeded via 273 Ds which decayed by emission of a 9.73 MeV alpha particle with a lifetime of 170 ms. This would have been assigned to an isomeric level. This data could not be confirmed and thus this isotope is currently unknown or unconfirmed. 272
Ds
In the first attempt to synthesise element 110, a 10 ms SF activity was assigned to 272Ds in the reaction 232Th(44Ca,4n). Given current understanding regarding stability, this isotope has been retracted from the Table of Isotopes.
8
Darmstadtium
9
Theoretical calculations on decay characteristics Theoretical calculation in a quantum tunneling model reproduces the experimental alpha decay half live data.[16] [17] It also predicts that the isotope 294110 would have alpha decay half life of the order of 311 years.[18] [19]
Fission of compound nuclei with Z=110 Experiments have been performed in 2004 at the Flerov Laboratory of Nuclear Reactions in Dubna studying the fission characteristics of the compound nucleus 280Ds. The nuclear reaction used is 232Th+48Ca. The result revealed how nuclei such as this fission predominantly by expelling closed shell nuclei such as 132Sn (Z=50, N=82).[20]
Theoretical calculations on evaporation cross sections The below table contains various targets-projectile combinations for which calculations have provided estimates for cross section yields from various neutron evaporation channels. The channel with the highest expected yield is given. DNS = Di-nuclear system ; σ = cross section Target
Projectile
CN
Channel (product)
~σ
max
Model
Ref
208
Pb
64
Ni
272
1n (271Ds)
10 pb
DNS
[21]
232
Th
48
Ca
280
4n (276Ds)
0.2 pb
DNS
[22]
230
48
278
4n (274Ds)
1 pb
DNS
[22]
278
4n (274Ds)
2 pb
DNS
[22]
Th
238
U
Ca
40
Ar
Ds Ds Ds Ds
Future Experiments The team at the HIRFL, Lanzhou, China are planning to restudy the reaction 238U(40Ar,xn) after recent calculations indicated a measurable yield in the 4n evaporation channel, leading to the new nuclide 274Ds. At the FLNR, scientists will study the new reaction 226Ra(50Ti,xn) in order to compare the yield with that obtained using 48Ca projectiles in order to ascertain the viability of using 50 Ti projectiles in SHE synthesis.
See also • Island of stability
External links • • • •
Darmstadtium.com [23] WebElements.com - Darmstadtium [24] IUPAC: Element 110 is named darmstadtium Apsidium - darmstadtium [26]
[25]
Darmstadtium
References [1] Prof S.Hofmann (private communication) [2] "Production and decay of 269110" (http:/ / www. springerlink. com/ content/ j51g8342h311x460/ ), Hofmann et al., Z. Phys. A., 1995, 350, 4. Retrieved on 2008-03-02 [3] "New elements-approaching Z=114" (http:/ / www. iop. org/ EJ/ abstract/ 0034-4885/ 61/ 6/ 002), S.Hoffman, Rep. Prog. Phys.,1998, 61, 639-689. Retrieved on 2008-03-02 [4] http:/ / www. iupac. org/ publications/ pac/ 2001/ pdf/ 7306x0959. pdf [5] ununnilium - Definitions from Dictionary.com (http:/ / dictionary. reference. com/ browse/ ununnilium) [6] "Confirmation of production of element 110 by the 208Pb(64Ni,n) reaction" (http:/ / prola. aps. org/ abstract/ PRC/ v67/ i6/ e064609), Ginter et al., Phys. Rev. C, 67, 064609 (2003). Retrieved on 2008-03-02 [7] "Confirmation of production of element 110 by the 208Pb(64Ni,n) reaction" (http:/ / repositories. cdlib. org/ cgi/ viewcontent. cgi?article=5446& context=lbnl), Ginter et al., LBNL repositories. Retrieved on 2008-03-02 [8] "Development of an Odd-Z-Projectile Reaction for Heavy Element Synthesis: 208Pb(64Ni,n)271Ds and 208 Pb(65Cu,n)272111" (http:/ / prola. aps. org/ abstract/ PRL/ v93/ i21/ e212702), Folden et al., Phys. Rev. Lett., 93, 212702 (2004). Retrieved on 2008-03-02 [9] "Development of an Odd-Z-Projectile Reaction for Heavy Element Synthesis: 208Pb(64Ni,n)271Ds and 208 Pb(65Cu,n)272111" (http:/ / repositories. cdlib. org/ cgi/ viewcontent. cgi?article=2704& context=lbnl), Folden et al., LBNL repositories. Retrieved on 2008-03-02 [10] "Production and decay of the isotope 271Ds (Z = 110)" (http:/ / www. springerlink. com/ content/ 74kn9nq0l28lmlge), Morita et al., Eur. Phys. J. A., 2004, 21, 2. Retrieved on 2008-03-02/ [11] "The new isotope 270110 and its decay products 266Hs and 262Sg" (http:/ / www. dnp. fmph. uniba. sk/ etext/ 68/ text/ Hofmann_et_al_EPJ_A10_(2001)_5. pdf), Hofmann et al., Eur. Phys. J. A., 10, 5-10 (2001). Retrieved on 2008-03-02 [12] "Evidence for the possible synthesis of element 110 produced by the 59Co+209Bi reaction" (http:/ / prola. aps. org/ abstract/ PRC/ v51/ i5/ pR2293_1), Ghiorso et al., Phys. Rev. C, 51, R2293-R2297 (1995). Retrieved on 2008-03-02 [13] "α decay of 273110: Shell closure at N=162" (http:/ / prola. aps. org/ abstract/ PRC/ v54/ i2/ p620_1), Lazarev et al., Phys. Rev. C, 54, 620-625 (1996). Retrieved on 2008-03-02 [14] see ununquadium [15] see ununoctium [16] P. Roy Chowdhury, C. Samanta, and D. N. Basu (26 January 2006). " α decay half-lives of new superheavy elements (http:/ / scitation. aip. org/ getabs/ servlet/ GetabsServlet?prog=normal& id=PRVCAN000073000001014612000001& idtype=cvips& gifs=yes)". Phys. Rev. C 73: 014612. doi: 10.1103/PhysRevC.73.014612 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 73. 014612). . [17] C. Samanta, P. Roy Chowdhury and D.N. Basu (2007). " Predictions of alpha decay half lives of heavy and superheavy elements (http:/ / www. sciencedirect. com/ science?_ob=ArticleURL& _udi=B6TVB-4NF4F0Y-2& _user=2806701& _rdoc=1& _fmt=& _orig=search& _sort=d& view=c& _acct=C000058844& _version=1& _urlVersion=0& _userid=2806701& md5=3f680654b5659191d67f31681a4cfc83)". Nucl. Phys. A 789: 142–154. doi: 10.1016/j.nuclphysa.2007.04.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2007. 04. 001). . [18] P. Roy Chowdhury, C. Samanta, and D. N. Basu (2008). " Search for long lived heaviest nuclei beyond the valley of stability (http:/ / scitation. aip. org/ getabs/ servlet/ GetabsServlet?prog=normal& id=PRVCAN000077000004044603000001& idtype=cvips& gifs=yes)". Phys. Rev. C 77: 044603. doi: 10.1103/PhysRevC.77.044603 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 77. 044603). . [19] P. Roy Chowdhury, C. Samanta, and D. N. Basu (2008). " Nuclear half-lives for α -radioactivity of elements with 100 ≤ Z ≤ 130 (http:/ / www. sciencedirect. com/ science?_ob=ArticleURL& _udi=B6WBB-4S26JRX-1& _user=2806701& _coverDate=03/ 14/ 2008& _alid=740505626& _rdoc=6& _fmt=high& _orig=search& _cdi=6706& _sort=d& _docanchor=& view=c& _ct=211& _acct=C000058844& _version=1& _urlVersion=0& _userid=2806701& md5=dc85a3a8a2ac1faa38c3804f16f86c13)". At. Data & Nucl. Data Tables. . [20] [21] [22] [23] [24]
see Flerov lab annual report 2004 inclusive http:/ / www1. jinr. ru/ Reports/ Reports_eng_arh. html http:/ / arxiv. org/ pdf/ 0707. 2588 http:/ / arxiv. org/ pdf/ 0803. 1117 http:/ / www. darmstadtium. com http:/ / www. webelements. com/ webelements/ elements/ text/ Ds/ index. html
[25] http:/ / www. iupac. org/ news/ archives/ 2003/ naming110. html [26] http:/ / www. apsidium. com/ elements/ 110. htm
10
Article Sources and Contributors
Article Sources and Contributors Darmstadtium Source: http://en.wikipedia.org/w/index.php?oldid=309171166 Contributors: Ahoerstemeier, AlimanRuna, Andres, Anger22, Beetstra, Blackoutjr, BlueEarth, Bobbob155, Borislav Dopudja, BrittanyP!, Bryan Derksen, CUSENZA Mario, Cameron Nedland, Carbuncle, CatherineMunro, Chris 73, Confederateman08, Cytokid101, DOSGuy, Danny, Darrien, David Gerard, Devleenasamanta, DopefishJustin, Drjezza, EPO, Eclecticology, Edgar181, El C, Emperorbma, Eras-mus, Eric119, Femto, Fonzy, Fpbecker, Gadfium, Gellersen, Greatpatton, Hak-kâ-ngìn, Halobeast, Hh, Hokanomono, Icairns, Ingolfson, JPisano, Jake Wartenberg, Joanjoc, John, Kachyna, Kaihsu, Kelovy, Kingdon, Ksbrown, Kurykh, Kwamikagami, Mav, Menchi, Mortdefides, Mysid, Neitherday, Nergaal, Noisy, Nozzleman, Olessi, Oliver Pereira, Pascal666, PiMaster3, PierreAbbat, Pol098, PookeyMaster, Poolkris, Pras, Quiddity, Remember, Rfc1394, Rifleman 82, Romanskolduns, Rursus, Sakus, Saperaud, Sceptre, Schneelocke, Scott3, Shell Kinney, Shogunzhu, Sionus, SirJibby, Sl, Smokefoot, Steffen1983, Stifynsemons, Stormwriter, Tagishsimon, Tetracube, The Anome, The Epopt, Timmy12, Tomchiukc, Try her philosophy, Urhixidur, Vicki Rosenzweig, Victor, Walkerma, Yekrats, 89 anonymous edits
Image Sources, Licenses and Contributors image:Ds-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Ds-TableImage.png License: GNU Free Documentation License Contributors: Dodo, Kwamikagami, Paddy, Saperaud File:Speaker Icon.svg Source: http://en.wikipedia.org/w/index.php?title=File:Speaker_Icon.svg License: Public Domain Contributors: Blast, G.Hagedorn, Mobius, 2 anonymous edits Image:270Ds decay scheme.png Source: http://en.wikipedia.org/w/index.php?title=File:270Ds_decay_scheme.png License: Public Domain Contributors: User:Drjezza
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11
Roentgenium
1
Roentgenium darmstadtium ← roentgenium → ununbium Au ↑ Rg ↓ (Uhu) Periodic Table Extended Periodic Table General Name, symbol, number
roentgenium, Rg, 111
Element category
transition metals
Group, Period, Block
11, 7, d
Appearance
unknown Standard atomic weight
−1
[280] g·mol
Electron configuration
14
predicted [Rn] 5f
Electrons per shell
9
2
6d 7s
2, 8, 18, 32, 32, 17, 2
Phase
presumably a solid CAS registry number
54386-24-2 Most-stable isotopes
Main article: Isotopes of roentgenium iso
NA
half-life
283
syn
10 min (est.)
282
syn
4 min (est.)
281
syn
1 min (est.)
280
syn
3.6 s
279
syn
170 ms
278
syn
274 272
Rg Rg Rg Rg Rg Rg Rg Rg
DM
α
DE (MeV)
DP
9.75
276
α
10.37
275
4.2 ms
α
10.69
274
syn
15 ms
α
11.23
270
syn
1.6 ms
α
11.02,10.82
268
References
Mt Mt Mt Mt Mt
Roentgenium
2
[1] Roentgenium (pronounced /rʌntˈgɛniəm/ ( listen)) is a chemical element in the periodic table that has the symbol Rg and atomic number 111.
Roentgenium is a radioactive synthetic element whose most stable known isotope has a mass number of 280 and a half-life of a 3.6 s.
Official discovery Roentgenium was officially discovered by Peter Armbruster, Gottfried Münzenberg, and their team working at the Gesellschaft für Schwerionenforschung (GSI) in Darmstadt, Germany on December 8, 1994.[2] Only three atoms of it were observed (all 272Rg), by the cold fusion between Nickel ions and a Bismuth target in a linear accelerator: 20983Bi
+
6428Ni
→
272111Rg
+
10n
In 2001, the IUPAC/IUPAP Joint Working Party (JWP) concluded that there was insufficient evidence for the discovery at that moment in time.[3] The GSI team repeated their experiment in 2000 and detected a further 3 atoms.[4] [5] In their 2003 report, the JWP decided that the GSI team should be acknowledged as the discoverers.[6]
Naming The name roentgenium (Rg) was proposed by the GSI team[7] in honor of the German physicist Wilhelm Conrad Röntgen, and was accepted as a permanent name on November 1, 2004.[8] Previously the element was known under the temporary IUPAC systematic element name unununium, Uuu.
Relativistic electronic structure The stable group 11 elements, copper, silver, and gold all have an outer electron configuration nd10(n+1)s1. For each of these elements, their first excited state has a configuration nd9(n+1)s2. Due to spin-orbit coupling between the s electrons, this state is split into a pair of energy levels. For copper, the difference in energy between the ground state and lowest excited state causes the metal to appear reddish. For silver, the energy gap widens and it become silvery. However, as Z increases, the excited levels are stabilised by relativistic effects and in gold the energy gap decreases again and it appears gold. For roentgenium, calculations indicate that the 6d97s2 level is stabilised to such an extent that it becomes the ground state. The resulting energy difference between the new ground state and the first excited state is similar to that of silver and roentgenium is expected to be silvery in appearance.[9]
Extrapolated chemical properties of eka-gold Oxidation states Element 111 is projected to be the ninth member of the 6d series of transition metals and the heaviest member of group 11 (IB) in the Periodic Table, below copper, silver, and gold. Each of the members of this group show different stable states. Copper forms a stable +2 state, whilst silver is predominantly found as silver(I) and gold as gold(III). Copper(I) and silver(II) are also relatively well-known. Roentgenium is therefore expected to predominantly form a stable +3 state.
Roentgenium
3
Chemistry The heavier members of this group are well known for their lack of reactivity or noble character. Silver and gold are both inert to oxygen. They are both however attacked by the halogens. In addition, silver is attacked by sulfur and hydrogen sulfide, highlighting its higher reactivity compared to gold. Roentgenium is expected to be even more noble than gold and can be expected to be inert to oxygen and halogens. The most-likely reaction is with fluorine to form a trifluoride, RgF3.
History of synthesis of isotopes by cold fusion This section deals with the synthesis of nuclei of roentgenium by so-called "cold" fusion reactions. These are processes which create compound nuclei at low excitation energy (~10-20 MeV, hence "cold"), leading to a higher probability of survival from fission. The excited nucleus then decays to the ground state via the emission of one or two neutrons only. 209
Bi(64Ni,xn)273−xRg (x=1)
First experiments to synthesize element 111 were performed by the Dubna team in 1986 using this cold fusion reaction. No atoms were identified that could be assigned to atoms of element 111 and a production cross-section limit of 4 pb was determined. After an upgrade of their facilities, the team at GSI successfully detected 3 atoms of 272Rg in their discovery experiment.[2] A further 3 atoms were synthesized in 2000.[4] The discovery of roentgenium was confirmed in 2003 when a team at RIKEN measured the decays of 14 atoms of 272Rg during the measurement of the 1n excitation function.[10] 208
Pb(65Cu,xn)273−xRg (x=1)
In 2004, as part of their study of odd-Z projectiles in cold fusion reactions, the team at LBNL detected a single atom of 272Rg in this new reaction.[11] [12]
History of synthesis of isotopes as decay products Isotopes of roentgenium have also been observed in the decay of heavier elements. Observations to date are outlined in the table below: Evaporation residue
Observed Rg isotope
288
280
[13]
287
279
[13]
282
278
[14]
278
274
[14]
115 115 113 113
Rg Rg Rg Rg
Roentgenium
4
Chronology of isotope discovery Isotope
Year discovered
Discoverer reaction
272
1994
273
unknown
274
2004
275
unknown
276
unknown
277
unknown
278
2006
237
[14]
279
2003
243
[13]
280
2003
243
[13]
Rg Rg Rg Rg Rg Rg Rg Rg Rg
209
Bi(64Ni,n)
209
Bi(70Zn,n)
[14]
Np(48Ca,3n) Am(48Ca,4n) Am(48Ca,3n)
Chemical yields of isotopes Cold fusion The table below provides cross-sections and excitation energies for cold fusion reactions producing roentgenium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
CN
1n
64
209
273
3.5 pb, 12.5 MeV
65
208
273
1.7 pb, 13.2 MeV
Ni
Bi
Cu
Pb
Rg Rg
2n
3n
Isotopes Five isotopes of roentgenium are known. The longest-lived of these is 280Rg, which decays through alpha decay and has a halflife of 3.6 seconds. The shortest-lived isotope is 272Rg, which decays through alpha decay and has a halflife of 1.6 ms.
Isomerism in roentgenium nuclides 274
Rg
Two atoms of 274Rg have been observed in the decay chains starting with 278Uut. The two events occur with different energies and with different lifetimes. In addition, the two entire decay chains appear to be different. This suggests the presence of two isomeric levels but further research is required.
Roentgenium
5
272
Rg
The direct production of 272Rg has provided four alpha lines at 11.37, 11.03, 10.82, and 10.40 MeV. The GSI measured a half-life of 1.6 ms whilst recent data from RIKEN has given a half-life of 3.8 ms. The conflicting data may be due to isomeric levels but the current data are insufficient to come to any firm assignments.
See also • Island of stability
External links • • • • • •
WebElements.com: Roentgenium [15] IUPAC: Proposal of name roentgenium for element 111 IUPAC: Element 111 is named roentgenium [17] Apsidium: Roentgenium 111 [18] Unununium Element At Chemicalelements.com [19] RSC.org: Roentgenium [20]
[16]
References [1] Prof S.Hofmann (private communication) [2] Hofmann, S. (1995). "The new element 111". Zeitschrift für Physik a Hadrons and Nuclei 350: 281. doi: 10.1007/BF01291182 (http:/ / dx. doi. org/ 10. 1007/ BF01291182). [3] Karol et al. (2001). " On the discovery of the elements 110–112 (http:/ / iupac. org/ publications/ pac/ 2001/ pdf/ 7306x0959. pdf)". Pure Appl. Chem. 73 (6): 959–967. doi: 10.1351/pac200173060959 (http:/ / dx. doi. org/ 10. 1351/ pac200173060959). . [4] Hofmann, S. (2002). "New results on elements 111 and 112". The European Physical Journal A 14: 147. doi: 10.1140/epja/i2001-10119-x (http:/ / dx. doi. org/ 10. 1140/ epja/ i2001-10119-x). [5] Hofmann et al.. " New results on element 111 and 112 (http:/ / www. gsi. de/ informationen/ wti/ library/ scientificreport2000/ Nuc_St/ 7/ ar-2000-z111-z112. pdf)". GSI report 2000. . Retrieved 2008-03-02. [6] " Karol et al. (http:/ / iupac. org/ publications/ pac/ 2003/ pdf/ 7510x1601. pdf)". Pure Appl. Chem. 75 (10): 1601–1611. 2003. . [7] Corish et al.. " Name and symbol of the element with atomic number 111 (http:/ / iupac. org/ reports/ provisional/ abstract04/ Corish_pr111. pdf)". IUPAC Provisional Recommendations. . Retrieved 2008-03-02. [8] Corish et al. (2004). " Name and symbol of the element with atomic number 111 (http:/ / iupac. org/ publications/ pac/ 2004/ pdf/ 7612x2101. pdf)". Pure Appl. Chem. 76 (12): 2101–2103. doi: 10.1351/pac200476122101 (http:/ / dx. doi. org/ 10. 1351/ pac200476122101). . [9] Turler, A. (2004). " Gas Phase Chemistry of Superheavy Elements (http:/ / wwwsoc. nii. ac. jp/ jnrs/ paper/ JN52/ j052Turler. pdf)". Journal of Nuclear and Radiochemical Sciences 5 (2): R19–R25. . [10] Morita, K (2004). "Status of heavy element research using GARIS at RIKEN". Nuclear Physics A 734: 101. doi: 10.1016/j.nuclphysa.2004.01.019 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2004. 01. 019). [11] Folden, C. M. (2004). "Development of an Odd-Z-Projectile Reaction for Heavy Element Synthesis: ^{208}Pb(^{64}Ni,n)^{271}Ds and ^{208}Pb(^{65}Cu,n)^{272}111". Physical Review Letters 93: 212702. doi: 10.1103/PhysRevLett.93.212702 (http:/ / dx. doi. org/ 10. 1103/ PhysRevLett. 93. 212702). [12] "Development of an Odd-Z-Projectile Reaction for Heavy Element Synthesis: 208Pb(64Ni,n)271Ds and 208 Pb(65Cu,n)272111" (http:/ / repositories. cdlib. org/ cgi/ viewcontent. cgi?article=2704& context=lbnl), Folden et al., LBNL repositories. Retrieved on 2008-03-02 [13] see ununpentium for details [14] see ununtrium for details [15] http:/ / webelements. com/ webelements/ elements/ text/ Rg/ index. html [16] http:/ / iupac. org/ reports/ provisional/ abstract04/ corish_311004. html [17] http:/ / iupac. org/ news/ archives/ 2004/ naming111. html [18] http:/ / www. apsidium. com/ elements/ 111. htm [19] http:/ / www. chemicalelements. com/ elements/ uuu. html
Roentgenium [20] http:/ / www. rsc. org/ chemsoc/ visualelements/ pages/ roentgenium. html
6
Article Sources and Contributors
Article Sources and Contributors Roentgenium Source: http://en.wikipedia.org/w/index.php?oldid=306270507 Contributors: 1337sword, AAP52, ABF, Ahoerstemeier, AlimanRuna, Allison Connors, Altenmann, AmateurHistorian, Andres, Andros 1337, AySz88, Badgermont, Beetstra, Benjamintchip, BigT27, Bilbo pingouin, Bkell, BlueEarth, Bogdangiusca, Boo516, Bryan Derksen, Cameron Nedland, Captain-tucker, Chris 73, Ctachme, DOSGuy, DabMachine, Darrien, Dblecros, Djr32, DopefishJustin, Dr Debug, Drjezza, Dungodung, ESkog, Edgar181, Emperorbma, Eoghanacht, Eric119, Excirial, Fabrictramp, Femto, Feneeth of Borg, Fonzy, Foobar, Frankhollymaya, Furocumarine, Greenguy746, Gwernol, Haham hanuka, Hak-kâ-ngìn, Hede2000, Heron, Hobomonkey92, IRP, Icairns, InFairness, JForget, Jaraalbe, Jni, Joffan, John, Jolb, Kelovy, King of Hearts, Kingdon, Kurykh, Kwamikagami, Kwekubo, Larskris, MER-C, Maedin, Materialscientist, MathStuf, Mav, Mbenzdabest, Mhking, Michael Zimmermann, Minghong, Mxn, N-true, NapoliRoma, Natalie Erin, NawlinWiki, Nergaal, Nihiltres, Noisy, Nonagonal Spider, Nozzleman, Ojigiri, Olin, Ortolan88, Philip Trueman, Polly, Poolkris, Pouya, Project2501a, Rares brodeanu, Redvers, Remember, Rknasc, Roentgenium111, Romanskolduns, RoryReloaded, Rursus, Saperaud, Schneelocke, Scyphiform, Sengkang, Shimgray, Sidonuke, SkE, SkyLined, Sl, Snowdog, SomeKindOfCamel, Sportzplyr9090, Stone, Tagishsimon, Tarry, TedE, Tetracube, The Obento Musubi, The Ronin, TheAmber, Timc, Tlogmer, Tobyc75, Tomchiukc, TwoOneTwo, Typhlosion, Ukexpat, Until It Sleeps, Vanisheduser12345, Versus22, Vezhlys, Vuo, Whkoh, WhoopeeDoo, X!, 180 anonymous edits
Image Sources, Licenses and Contributors image:Rg-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Rg-TableImage.png License: GNU Free Documentation License Contributors: Conscious, Dbenbenn, Kwamikagami, Paddy, Saperaud, Thuresson, 2 anonymous edits File:Speaker Icon.svg Source: http://en.wikipedia.org/w/index.php?title=File:Speaker_Icon.svg License: Public Domain Contributors: Blast, G.Hagedorn, Mobius, 2 anonymous edits
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7
Ununbium
1
Ununbium roentgenium ← ununbium → ununtrium Hg ↑ Uub ↓ (Uhb) Periodic Table Extended Periodic Table General Name, symbol, number
ununbium, Uub, 112
Element category
transition metals
Group, Period, Block
12, 7, d
Appearance
unknown Standard atomic weight
−1
[285] g·mol
Electron configuration
14
[Rn] 5f
Electrons per shell
10
6d
2
7s
2, 8, 18, 32, 32, 18, 2
Phase
unknown CAS registry number
54084-26-3 Most-stable isotopes
Main article: Isotopes of ununbium iso
NA
half-life
DM
DE (MeV)
DP
syn
29 s
α
9.15,9.03?
281
285b
syn
8.9 m
α
8.63
281b
284
syn
97 ms
SF
283
syn
4s
~80% α
9.53,9.32,8.94
279
11.45,11.32
273
285
Uub Uub ?
Uub Uub
Ds Ds ?
Ds
~20% SF 283b
syn
~ 7.0 m
100% SF
282
syn
0.8 ms
SF
277
syn
0.7 ms
α
Uub ??
Uub Uub
Ds
References
Ununbium is a chemical element in the periodic table that has the temporary symbol Uub and the atomic number 112. "Ununbium" (commonly pronounced /juːˈnʌnbiəm/,[1] officially
Ununbium
2 [2]
English pronunciation: /uːˈnuːnbiəm/ (
listen)
) is a IUPAC systematic element name, used
until the element gets an accepted name. Ununbium was first created by the GSI in 1996, who have proposed the permanent name [3] copernicium English pronunciation: /koʊpərˈniːsiəm/ ( listen) and the symbol Cp.[4] It is the highest-numbered superheavy element to be officially recognised by IUPAC. This name is expected to be officially endorsed by IUPAC in January 2010, after six months for discussion.[5] The most stable isotope discovered to date is 285Uub with a half-life of ~30 s. In total, about 75 atoms of ununbium have been detected using various nuclear reactions.[6] An unconfirmed isotope, 285bUub, has a possible half-life of ~9 minutes, and would be one of the longest-lived superheavy isotopes known to date. Recent experiments strongly suggest that ununbium behaves as a typical member of group 12, demonstrating properties consistent with a volatile metal.[7]
Official discovery Ununbium was first created on February 9, 1996 at the Gesellschaft für Schwerionenforschung (GSI) in Darmstadt, Germany by Sigurd Hofmann, Victor Ninov et al.[8] This element was created by firing accelerated zinc-70 nuclei at a target made of lead-208 nuclei in a heavy ion accelerator. A single atom (the second has subsequently been dismissed) of ununbium was produced with a mass number of 277.[8] 20882Pb
+
7030Zn
→
278112Uub
→
277112Uub
+
10n
In May 2000, the GSI successfully repeated the experiment to synthesise a further atom of Uub-277.[9] [10] This reaction was repeated at RIKEN using the GARIS set-up in 2004 to synthesise two further atoms and confirm the decay data reported by the GSI team.[11] The IUPAC/IUPAP Joint Working Party (JWP) assessed the claim of discovery by the GSI team in 2001[12] and 2003.[13] In both cases, they found that there was insufficient evidence to support their claim. This was primarily related to the contradicting decay data for the known isotope 261Rf. However, between 2001-2005, the GSI team studied the reaction 248 Cm(26Mg,5n)269Hs, and were able to confirm the decay data for 269Hs and 261Rf. It was found that the existing data on 261Rf was for an isomer,[14] now designated 261a Rf. In May 2009, the JWP reported on the claims of discovery of element 112 again and officially recognised the GSI team as the discoverers of element 112.[15] This decision was based on recent confirmation of the decay properties of daughter nuclei as well as the confirmatory experiments at RIKEN.[16]
Naming The element with Z=112 is historically known as eka-mercury. Ununbium is a temporary IUPAC systematic element name. Research scientists usually refer to the element simply as element 112 (or just E112). After acknowledging their discovery, the IUPAC has asked the discovery team at GSI to suggest a permanent name for ununbium.[17] On the 14th July 2009, they proposed "copernicium" with the element symbol "Cp", after Nicolaus Copernicus "to honor an outstanding scientist, who changed our view of the world."[18] IUPAC has not yet officially recognized this name, pending the results of a six month discussion period among the scientific community.[5]
Ununbium Some news articles have referred to the suggested name as "copernicum" in error.[19] However, the IUPAC only allows the suffix -ium for new elements. Alternative spellings have been suggested to Hofmann, namely "copernicum", "copernium", and "kopernikium" (Kp), and Hofmann has said that the team had discussed the possibility of "copernicum" or "kopernikum", but that they had agreed on "copernicium" in order to comply with current IUPAC rules.[20]
Extrapolated chemical properties Oxidation states Element 112 is projected to be the last member of the 6d series of transition metals and the heaviest member of group 12 (IIB) in the Periodic Table, below zinc, cadmium and mercury. Each of the members of this group show a stable +2 oxidation state. In addition, mercury(I), Hg2+2, is also well known. Element 112 is therefore expected to form a stable +2 state.
Chemistry The known members of group 12 all react with oxygen and sulfur directly to form the oxides and sulfides, MO and MS, respectively. Mercury(II) oxide, HgO, can be decomposed by heat to the liquid metal. Mercury also has a well known affinity for sulfur. Therefore, element 112 should form an analogous oxide UubO and sulfide UubS. In their halogen chemistry, all the metals form the ionic difluoride MF2 upon reaction with fluorine. The other halides are known but for mercury, the soft nature of the Hg(II) ion leads to a high degree of covalency and HgCl2, HgBr2 and HgI2 are low-melting, volatile solids. Therefore, element 112 is expected to form an ionic fluoride, UubF2, but volatile halides, UubCl2, UubBr2 and UubI2. In addition, mercury is well known for its alloying properties, with the concomitant formation of amalgams, especially with gold and silver. It is also a volatile metal and is monatomic in the vapour phase. Element 112 is therefore also predicted to be a volatile metal which readily combines with gold to form a Au-Uub metal-metal bond.
Experimental chemistry Atomic gas phase Ununbium is expected to have the ground state electron configuration [Rn]5f14 6d10 7s2 and thus belong to group 12 of the Periodic Table. As such, it should behave as the heavier homologue of mercury (Hg) and form strong binary compounds with noble metals like gold. Experiments probing the reactivity of ununbium have focused on the adsorption of atoms of element 112 onto a gold surface held at varying temperatures, in order to calculate an adsorption enthalpy. Due to possible relativistic stabilisation of the 7s electrons, leading to radon-like properties, experiments were performed with the simultaneous formation of mercury and radon radioisotopes, allowing a comparison of adsorption characteristics. The first experiments were conducted using the 238U(48Ca,3n)283112 reaction. Detection was by spontaneous fission of the claimed 5 min parent isotope. Analysis of the data indicated that ununbium was more volatile than mercury and had noble-gas properties. However, the confusion regarding the synthesis of 283112 has cast some doubt on these experimental results.
3
Ununbium Given this uncertainty, between April-May 2006 at the JINR, a FLNR-PSI team conducted experiments probing the synthesis of this isotope as a daughter in the nuclear reaction 242 Pu(48Ca,3n)287114. In this experiment, two atoms of 283112 were unambiguously identified and the adsorption properties indicated that ununbium is a more volatile homologue of mercury, due to formation of a weak metal-metal bond with gold, placing it firmly in group 12. In April 2007 this experiment was repeated and a further 3 atoms of 283112 were positively identified. The adsorption property was confirmed and indicated that element 112 has adsorption properties completely in agreement with being the heaviest member of group 12.[21]
History of synthesis of isotopes by cold fusion This section deals with the synthesis of nuclei of ununbium by so-called "cold" fusion reactions. These are processes which create compound nuclei at low excitation energy (~10-20 MeV, hence "cold"), leading to a higher probability of survival from fission. The excited nucleus then decays to the ground state via the emission of one or two neutrons only. 208
Pb(70Zn,xn)278-xUub (x=1)
The team at GSI first studied this reaction in 1996 and detected two decay chains of 277 Uub.[8] In a review of the data in 2000, the first decay chain was retracted. In a repeat of the reaction in 2000 they were able to synthesise a further atom. They attempted to measure the 1n excitation function in 2002 but suffered from a failure of the Zn-70 beam. The unofficial discovery of 277Uub was confirmed in 2004 at RIKEN who detected a further 2 atoms of the isotope and were able to confirm the decay data for the entire chain. 208
Pb(68Zn,xn)276-x112
Following the successful synthesis of 277Uub, the GSI team performed a reaction using a 68 Zn projectile in 1997 in an effort to study the effect of isospin (neutron richness) on the chemical yield. The experiment was initiated following the discovery of a yield enhancement during the synthesis of darmstadtium isotopes using 62Ni and 64Ni ions. No decay chains of 275112 were detected leading to a cross section limit of 1.2 pb. However, the revision of the yield for the 70Zn reaction to 0.5 pb does not rule out a similar yield for this reaction. 184
W(88Sr,xn)272-x112
In 1990, after some early indications for the formation of isotopes of element 112 in the irradiation of a tungsten target with multi-GeV protons, a collaboration between GSI and the University of Jerusalem studied the above reaction. They were able to detect some spontaneous fission activity and a 12.5 MeV alpha decay, both of which they tentatively assigned to the radiative capture product 272112 or the 1n evaporation residue 271112. Both the TWG and JWP have concluded that a lot more research is required to confirm these conclusions.
4
Ununbium
History of synthesis of isotopes by hot fusion This section deals with the synthesis of nuclei of ununbium by so-called "hot" fusion reactions. These are processes which create compound nuclei at high excitation energy (~40-50 MeV, hence "hot"), leading to a reduced probability of survival from fission and quasi-fission. The excited nucleus then decays to the ground state via the emission of 3-5 neutrons. Fusion reactions utilizing 48Ca nuclei usually produce compound nuclei with intermediate excitation energies (~30-35 MeV) and are sometimes referred to as "warm" fusion reactions. This leads, in part, to relatively high yields from these reactions. 238
U(48Ca,xn)286-x112 (x=3,4)
In 1998, the team at the Flerov Laboratory of Nuclear Research began a research program using Ca-48 nuclei in "warm" fusion reactions leading to superheavy elements (SHE's). In March 1998, they claimed to have synthesised the element (2 atoms) in this reaction. The product, 283Uub, had a claimed half-life of 5 min, decaying by spontaneous fission (SF).[22] The long lifetime of the product initiated first chemical experiments on the gas phase atomic chemistry of element 112. In 2000, Yuri Yukashev at Dubna repeated the experiment but was unable to observe any spontaneous fission from 5 min activities. The experiment was repeated in 2001 and an accumulation of 8 SF fragments were found in the low temperature section, indicating that ununbium had radon-like properties. However, there is now some serious doubt about the origin of these results. In order to confirm the synthesis, the reaction was successfully repeated by the same team in Jan 2003, confirming the decay mode and half life. They were also able to calculate an estimate of the mass of the SF activity to ~285 lending support to the assignment.[23] The team at LBNL entered the debate and performed the reaction in 2002. They were unable to detect any SF activities and calculated a cross section limit of 1.6 pb for the detection of a single event.[24] The reaction was repeated in 2003-2004 by the team at Dubna using a slightly different set-up, the Dubna Gas Filled Recoil Separator (DGFRS). This time, 283Uub was found to decay by emission of a 9.53 MeV alpha-particle with a half-life of 4 seconds. 282Uub was also observed in the 4n channel.[25] In 2003, the team at GSI entered the debate and performed a search for the 5 minute SF activity in chemical experiments. Like the Dubna team, they were able to detect 7 SF fragments in the low temperature section. However, these SF events were uncorrelated, suggesting they were not from actual direct SF of element 112 nuclei and raised doubts about the original indications for radon-like properties.[26] After the announcement from Dubna of different decay properties for 283112, the GSI team repeated the experiment in September 2004. They were unable to detect any SF events and calculated a cross section limit of ~ 1.6 pb for the detection of one event, not in contradiction with the reported 2.5 pb yield by Dubna. In May 2005, the GSI performed a physical experiment and identified a single atom of 283 112 decaying by SF with a short lifetime suggesting a previously unknown SF branch.[27] However, initial work by Dubna had detected several direct SF events but had assumed that the parent alpha decay had been missed. These results indicated that this was not the case.
5
Ununbium
6
In 2006, the new decay data on 283112 was confirmed by a joint PSI-FLNR experiment aimed at probing the chemical properties of ununbium. Two atoms of 283Uub were observed in the decay of the parent 287Uuq nuclei. The experiment indicated that contrary to previous experiments, element 112 behaves as a typical member of group 12, [7] demonstrating properties of a volatile metal. Finally, the team at GSI successfully repeated their physical experiment in Jan 2007 and detected 3 atoms of 283112, confirming both the alpha and SF decay modes. [28] As such, the 5 min SF activity is still unconfirmed and unidentified. It is possible that it refers to an isomer, namely 283bUub, whose yield is obviously dependent upon the exact production methods. 233
U(48Ca,xn)281-x112
The team at FLNR studied this reaction in 2004. They were unable to detect any atoms of element 112 and calculated a cross section limit of 600 fb. The team concluded that this indicated that the neutron mass number for the compound nucleus had an effect on the yield of evaporation residues. [25]
Synthesis of isotopes as decay products Element 112 has also been observed as decay products of elements 114, 116 and 118 (see ununoctium). Evaporation Residue
Observed Uub isotope
293
289
285
292
288
284
291
287
283
294
290
116 , 116 , 116 , 118 ,
114
112
114
112
114 116 ,
112
286
282
114
112
As an example, in May 2006, the Dubna team (JINR) identified the decay of ununoctium via the alpha decay sequence: 294118Uuo
→
290116Uuh
→
286114Uuq
→
282
Uub as a final product in
282112Uub
It was found that the final nucleus undergoes spontaneous fission.[29]
Chronology of isotope discovery Isotope
Year discovered
277
1996
278
unknown
279
unknown
280
unknown
281
unknown
282
2004
Uub Uub Uub Uub Uub Uub
discovery reaction 208
Pb(70Zn,n)
238
U(48Ca,4n)
Ununbium
7 2002
244
283b
1998
238
284
2002
244
285
1999
244
1999
244
283
Uub Uub ??
Uub Uub
285b
Uub ?
Pu(48Ca,5n) U(48Ca,3n) Pu(48Ca,4n) Pu(48Ca,3n) Pu(48Ca,3n)
Chemical yields of isotopes Cold fusion The table below provides cross-sections and excitation energies for cold fusion reactions producing ununbium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
CN
1n
70
208
278
0.5 pb , 10.0;12.0 MeV
68
208
276
< 1.2 pb , 11.3;12.8 MeV
Zn
Pb
Zn
Pb
Uub Uub
2n
3n
Hot fusion The table below provides cross-sections and excitation energies for hot fusion reactions producing ununbium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
CN
3n
48
238
286
2.5 pb , 35.0 MeV
48
233
281
< 0.6 pb , 34.9 MeV
Ca
U
Ca
U
Uub Uub
4n
5n
0.6 pb
Isomerism in ununbium nuclides 285a,b
112
In the synthesis of 289114 and 293116, a 8.63 MeV alpha-decaying activity has been detected with a half-life of 8.9 minutes. Although unconfirmed in recent experiments, it is highly possible that this is associated with an isomer, namely 285b112. 283a,b
112
First experiments on the synthesis of 283112 produced a SF activity with half-life ~5 min. This activity was also observed from the alpha decay of 287114. The decay mode and half-life were also confirmed in a repeat of the first experiment. However, more recently,283112 has been observed to undergo 9.52 MeV alpha decay and SF with a half-life of 3.9 s. These results suggest the assignment of the two activities to two different isomeric levels in 283112, creating 283a112 and 283b112. Further research is required to address these discrepancies.
Ununbium
8
Retracted isotopes 281
112
In the claimed synthesis of 293118 in 1999 (see ununoctium) the isotope 281112 was idenitified as decaying by emission of a 10.68 MeV alpha particle with half-life 0.90 ms. The claim was retracted in 2001 and hence this ununbium isotope is currently unknown or unconfirmed.
Fission of compound nuclei with Z=112 Several experiments have been performed between 2001-2004 at the Flerov Laboratory of Nuclear Reactions in Dubna studying the fission characteristics of the compound nucleus 286 112. The nuclear reaction used is 238U+48Ca. The results have revealed how nuclei such as this fission predominantly by expelling closed shell nuclei such as 132Sn (Z=50, N=82). It was also found that the yield for the fusion-fission pathway was similar between 48Ca and 58 Fe projectiles, indicating a possible future use of 58Fe projectiles in superheavy element formation.[30]
Theoretical calculations on evaporation cross sections The below table contains various targets-projectile combinations for which calculations have provided estimates for cross section yields from various neutron evaporation channels. The channel with the highest expected yield is given. DNS = Di-nuclear system ; σ = cross section Target
Projectile
CN
Channel (product)
~σ
max
Model
Ref
208
70
278
1n (277112)
1.5 pb
DNS
[31]
208
67
Zn
275
1n (274112)
2 pb
DNS
[31]
U
48
Ca
286
4n (282112)
0.2 pb
DNS
[32]
Pu
40
284
4n (280112)
0.1 pb
DNS
[32]
36
S
286
4n (282112)
5 pb
DNS
[32]
30
Si
282
3n (279112)
10 pb
DNS
[32]
Pb Pb
238
244
250
Cm
252
Cf
Zn
Ar
112 112 112 112 112 112
Ununbium
9
See also • Island of stability • Polonium
External links • WebElements.com: Copernicium
[33]
References [1] " ununbium (http:/ / reference. aol. com/ columbia/ _a/ ununbium/ 20051207161909990010)". Columbia Encyclopedia. . [2] http:/ / media. iupac. org/ publications/ pac/ 1979/ pdf/ 5102x0381. pdf [3] Prof. S. Hofmann (private communication) [4] " 'Copernicium' proposed as name for newly discovered element 112 (http:/ / www. physorg. com/ news166791177. html)", PhysOrg.com, 2009-07-14, , retrieved 2009-07-14 [5] New element named 'copernicium' (http:/ / news. bbc. co. uk/ 1/ hi/ sci/ tech/ 8153596. stm), BBC News, Thu 16 July 2009 [6] see references in this article relating to 277112, 282112 and 283112, as well as references in ununquadium, ununhexium and ununoctium regarding observed daughter nuclei [7] R. Eichler, et al. (2007). "Chemical Characterization of Element 112". Nature 447: 72–75. doi: 10.1038/nature05761 (http:/ / dx. doi. org/ 10. 1038/ nature05761). [8] S. Hofmann, et al. (1996). "The new element 112". Zeitschrift für Physik: A Hadrons and Nuclei 354 (1): 229–230. doi: 10.1007/BF02769517 (http:/ / dx. doi. org/ 10. 1007/ BF02769517). [9] Hofmann et al. (2002). "New Results on Element 111 and 112". European Physical Journal A Hadrons and Nuclei 14 (2): 147–57. doi: 10.1140/epja/i2001-10119-x (http:/ / dx. doi. org/ 10. 1140/ epja/ i2001-10119-x). [10] Hofmann et al. (2000). " New Results on Element 111 and 112 (http:/ / www. gsi. de/ informationen/ wti/ library/ scientificreport2000/ Nuc_St/ 7/ ar-2000-z111-z112. pdf)". GSI Scientific Report 2000. . [11] K. Morita (2004). "Decay of an Isotope 277112 produced by 208Pb + 70Zn reaction". Exotic Nuclei (EXON2004). Proceedings of the International Symposium. World Scientific. pp. 188-191. doi: 10.1142/9789812701749_0027 (http:/ / dx. doi. org/ 10. 1142/ 9789812701749_0027). [12] P. J. Karol; H. Nakahara; B. W. Petley; E. Vogt (2001). " On the Discovery of the Elements 110—112 (http:/ / www. iupac. org/ publications/ pac/ 2001/ pdf/ 7306x0959. pdf)" (IUPAC Technical Report). Pure Appl. Chem. 73 (6): 959–967. doi: 10.1351/pac200173060959 (http:/ / dx. doi. org/ 10. 1351/ pac200173060959). . [13] P. J. Karol; H. Nakahara; B. W. Petley; E. Vogt (2003). " On the Claims for Discovery of Elements 110, 111, 112, 114, 116 and 118 (http:/ / www. iupac. org/ publications/ pac/ 2003/ pdf/ 7510x1601. pdf)" (IUPAC Technical Report). Pure Appl. Chem. 75 (10): 1061–1611. doi: 10.1351/pac200375101601 (http:/ / dx. doi. org/ 10. 1351/ pac200375101601). . [14] R. Dressler; A. Türler (2001). " Evidence for Isomeric States in 261Rf (http:/ / lch. web. psi. ch/ pdf/ anrep01/ B-02heavies. pdf)". Annual Report 2001. Paul Scherrer Institute. . [15] http:/ / www. gsi. de/ portrait/ Pressemeldungen/ 10062009-1_e. html [16] R.C.Barber; H.W.Gaeggeler;P.J.Karol;H. Nakahara; E.Vardaci; E. Vogt (2009). " Discovery of the element with atomic number 112 (http:/ / media. iupac. org/ publications/ pac/ asap/ pdf/ PAC-REP-08-03-05. pdf)" (IUPAC Technical Report). Pure Appl. Chem.. doi: 10.1351/PAC-REP-08-03-05 (http:/ / dx. doi. org/ 10. 1351/ PAC-REP-08-03-05). . [17] New Chemical Element In The Periodic Table (http:/ / www. sciencedaily. com/ releases/ 2009/ 06/ 090611210039. htm), www.sciencedaily.com [18] [http://www.gsi.de/portrait/Pressemeldungen/14072009_e.html July 14, 2009 - Element 112 shall be named “copernicium”, www.gsi.de [19] http:/ / www. popsci. com/ scitech/ article/ 2009-07/ element-112-named-copernicum [20] private email from Hofman [21] H. W. Gäggeler (2007). " Gas Phase Chemistry of Superheavy Elements (http:/ / lch. web. psi. ch/ pdf/ TexasA& M/ TexasA& M. pdf)". Paul Scherrer Institute. . [22] Oganessian et al. (1999). "Search for new isotopes of element 112 by irradiation of 238U with 48Ca". Eur. Phys. J. A 5 (1): 63–68. doi: 10.1007/s100500050257 (http:/ / dx. doi. org/ 10. 1007/ s100500050257). [23] Yu Ts Oganessian et al. (2004). "Second Experiment at VASSILISSA separator on the synthesis of the element 112". Eur. Phys. J. A 19 (1): 3–6. doi: 10.1140/epja/i2003-10113-4 (http:/ / dx. doi. org/ 10. 1140/ epja/
Ununbium i2003-10113-4). [24] W. Loveland, K. E. Gregorich, J. B. Patin, D. Peterson, C. Rouki, P. M. Zielinski, and K. Aleklett (2002). "Search for the production of element 112 in the 48Ca+238U reaction". Phys. Rev. C 66 (4): 044617. doi: 10.1103/PhysRevC.66.044617 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 66. 044617). [25] Yu. Ts. Oganessian et al. (2004). "Measurements of cross sections and decay properties of the isotopes of 233,238 242 248 48 elements 112, 114, and 116 produced in the fusion reactions U, Pu , and Cm+ Ca"]". Phys. Rev. C 70: 064609. doi: 10.1103/PhysRevC.70.064609 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 70. 064609). [26] S. Soverna (2003). Indication for a gaseous element 112 (http:/ / www. gsi. de/ informationen/ wti/ library/ scientificreport2003/ files/ 167. pdf). 2003. GSI Scientific Report. pp. 187. . [27] S. Hofmann, et al. (2005). Search for Element 112 Using the Hot Fusion Reaction 48Ca + 238U (http:/ / www. gsi. de/ informationen/ wti/ library/ scientificreport2005/ PAPERS/ NUSTAR-SHE-PHYS-01. pdf). 2005. GSI Scientific Report. pp. 191. . [28] S. Hofmann et al. (2007). "The reaction 48Ca + 238U -> 286112* studied at the GSI-SHIP". Eur. Phys. J. A 32 (3): 251–260. doi: 10.1140/epja/i2007-10373-x (http:/ / dx. doi. org/ 10. 1140/ epja/ i2007-10373-x). [29] Oganessian, Yu. Ts.; et al. (2006-10-09). "Synthesis of the isotopes of elements 118 and 116 in the 249Cf and 245 Cm+48Ca fusion reactions". Physical Review C 74 (4): 044602. doi: 10.1103/PhysRevC.74.044602 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 74. 044602). [30] [31] [32] [33]
see Flerov lab annual reports 2001-2004 inclusive http:/ / www1. jinr. ru/ Reports/ Reports_eng_arh. html http:/ / arxiv. org/ pdf/ 0707. 2588 http:/ / arxiv. org/ pdf/ 0904. 2994 http:/ / www. webelements. com/ copernicium/
10
Article Sources and Contributors
Article Sources and Contributors Ununbium Source: http://en.wikipedia.org/w/index.php?oldid=308752456 Contributors: 04redsox07, 84user, Ahoerstemeier, AlimanRuna, Altenmann, Alway Arptonshay, Andre Engels, Arakunem, Bagel7, Beetstra, Bigguy03j, Blackcat0030, BlueEarth, BlueNovember, Bondegezou, Brett Dunbar, Bryan Derksen, CBDunkerson, CYD, Cacahueten, Carbuncle, Chris 73, Coemgenus, Cosmium, Cryptic C62, DMacks, Dagonweb, Dajwilkinson, Danny, Darrien, David Gerard, Devleenasamanta, Discospinster, Dor Cohen, Drini, Drjezza, Earthlyreason, Edgar181, Emperorbma, Eras-mus, Eric119, Eszett, Eu-151, Eugene van der Pijll, Femto, Ferengi, Flying Jazz, Fonzy, Free-mind, Fvw, Garden, Geo Swan, Georgia guy, Glenn L, Gveret Tered, Hu, Icairns, Iridescent, Jagged 85, John, JohnyDog, Kelovy, Kingdon, Kurykh, Kwamikagami, Loamawanafuna, Lumos3, Lusanaherandraton, Mahahahaneapneap, Materialscientist, Mathiasrex, Matt B., Matěj Grabovský, Mav, Minesweeper, Nergaal, Nightstallion, Nihiltres, Noisy, Nozzleman, Ortolan88, Pascal666, PierreAbbat, Plau, Poolkris, Potatoswatter, Prnkap, Propaniac, QYV, Rcsheets, RedAndr, Redrose64, Remember, Reyk, Rhoppenrath-, Rich Farmbrough, RickK, Rifleman 82, Roentgenium111, Romanskolduns, Rursus, Saperaud, Schneelocke, Sesshomaru, Shimgray, SimonP, Siroxo, Sl, Stevey7788, Stismail, Stone, StuffOfInterest, Sugarfish, Sugarpine, Svante, Tagishsimon, Teetaweepo, Tetracube, Textangel, The Hokkaido Crow, The wub, Tim Q. Wells, Timc, Timwi, Trollminator, Until It Sleeps, Vicki Rosenzweig, Vuo, Vyznev Xnebara, Waldir, Warut, Wildthing61476, Xmrbearx, YeshuaDavid, 133 anonymous edits
Image Sources, Licenses and Contributors image:Uub-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Uub-TableImage.png License: GNU Free Documentation License Contributors: Daniel Mayer and Arnaud Gaillard File:Speaker Icon.svg Source: http://en.wikipedia.org/w/index.php?title=File:Speaker_Icon.svg License: Public Domain Contributors: Blast, G.Hagedorn, Mobius, 2 anonymous edits
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
11
Ununtrium
1
Ununtrium ununbium ← ununtrium → ununquadium Tl ↑ Uut ↓ (Uht) Periodic Table Extended Periodic Table General Name, symbol, number
ununtrium, Uut, 113
Element category
presumably poor metals
Group, Period, Block
13, 7, p
Appearance
possibly light or dark element Standard atomic weight
−1
[284] g·mol
Electron configuration
14
10
2
1
perhaps [Rn] 5f 6d 7s 7p (guess based on thallium)
Electrons per shell
2, 8, 18, 32, 32, 18, 3
Phase
presumably a solid CAS registry number
54084-70-7 Most-stable isotopes
Main article: Isotopes of ununtrium iso
NA
half-life
284
syn
0.49 s
283
syn
0.10 s
282
syn
278
syn
Uut Uut Uut Uut
DM α
DE (MeV)
DP
10.00
280
α
10.12
279
73 ms
α
10.63
278
0.34 ms
α
11.68
274
Rg Rg Rg Rg
References
Ununtrium (pronounced /juːˈnʌntriəm/; officially, the two initial u's are to be pronounced [1] English pronunciation: /uː/ ( listen) ) is the temporary name of a synthetic element in the periodic table that has the temporary symbol Uut and has the atomic number 113. It has been synthesised directly in both "cold" and "warm" fusion reactions. It was first observed in the decay of ununpentium. Only eight atoms of ununtrium have been observed to date. Following periodic trends it is expected to be a soft, silvery metal.
Ununtrium
2
Discovery profile The first report of ununtrium was in August 2003 when it was identified as a decay product of ununpentium. These results were published on February 1, 2004, by a team composed of Russian scientists at Dubna (Joint Institute for Nuclear Research), and American scientists at the Lawrence Livermore National Laboratory.[2] [3] On July 23, 2004, a team of Japanese scientists at RIKEN detected a single atom of 278Uut using the cold fusion reaction between Bismuth-209 and zinc-70. They published their results on September 28, 2004.[4]
Support for their claim appeared in 2004 when scientists at the Institute of Modern Physics (IMP) identified 266Bh as decaying with identical properties to their single event (see bohrium). The RIKEN team produced a further atom on April 2, 2005, although the decay data was different from the first chain, and may be due to the formation of a meta-stable isomer. The Dubna-Livermore collaboration has strengthened their claim for the discovery of ununtrium by conducting chemical experiments on the decay daughter 268Db. In experiments in Jun 2004 and Dec 2005, the Dubnium isotope was successfully identified by milking the Db fraction and measuring any SF activities. Both the half-life and decay mode were confirmed for the proposed 268Db which lends support to the assignment of Z=115 and Z=113 to the parent and daughter nuclei.[5] [6] Theoretical estimates of alpha-decay half-lives of alpha-decay chains from element 113 are in good agreement with the experimental data.[7]
Naming The element with Z=113 is historically known as eka-thallium. Ununtrium (Uut) is a temporary IUPAC systematic element name. Research scientists usually refer to the element simply as element 113 (E113).
Proposed names by claimants Claims to the discovery of element 113 have been put forward by Dmitriev of the Dubna team and Morita of the RIKEN team. The JWP will decide to whom the right to suggest a name will be given. The IUPAC have the final say on the official adoption of a name. The table below gives the names that the teams above have suggested and which can be verified by press interviews. Group
Proposed Name
Derivation
[8]
Japan - country of group claimants
[8]
RIKEN - institute of group claimants
RIKEN
Japonium
RIKEN
Rikenium
Ununtrium
Extrapolated chemical properties of eka-thallium Oxidation states Element 113 is projected to be the first member of the 7p series of non-metals and the heaviest member of group 13 (IIIA) in the Periodic Table, below thallium. Each of the members of this group show the group oxidation state of +III. However, thallium has a tendency to form only a stable +I state due to the "inert pair effect", explained by the relativistic stabilisation of the 7s-orbitals, resulting in a higher ionisation potential and weaker tendency to participate in bonding.
Chemistry Element 113 should portray eka-thallium chemical properties and should therefore form a monoxide, Uut2O, and monohalides, UutF, UutCl, UutBr and UutI. If the +III state is accessible, it is likely that it is only possible in the oxide, Uut2O3, and fluoride, UutF3.
History of synthesis of isotopes by cold fusion This section deals with the synthesis of nuclei of ununtrium by so-called "cold" fusion reactions. These are processes which create compound nuclei at low excitation energy (~10-20 MeV, hence "cold"), leading to a higher probability of survival from fission. The excited nucleus then decays to the ground state via the emission of one or two neutrons only. 209
Bi(70Zn,xn)279-x113 (x=1)
The synthesis of element 113 was first attempted in 1998 by the team at GSI using the above cold fusion reaction. In two separate runs, they were unable to detect any atoms and calculated a cross section limit of 900 fb.[9] They repeated the experiment in 2003 and lowered the limit further to 400 fb.[9] In late 2003, the emerging team at RIKEN using their efficient apparatus GARIS attempted the reaction and reached a limit of 140 fb. In December 2003-August 2004, they resorted to 'brute force' and performed an eight-month-long irradiation in which they increased the sensitivity to 51 fb. They were able to detect a single atom of 278113.[4] They repeated the reaction in several runs in 2005 and were able to synthesize a second atom. They calculated a record-low 31 fb for the cross section for the 2 atoms.
History of synthesis of isotopes by hot fusion This section deals with the synthesis of nuclei of ununtrium by so-called "hot" fusion reactions. These are processes which create compound nuclei at high excitation energy (~40-50 MeV, hence "hot"), leading to a reduced probability of survival from fission. The excited nucleus then decays to the ground state via the emission of 3-5 neutrons. Fusion reactions utilizing 48Ca nuclei usually produce compound nuclei with intermediate excitation energies (~30-35 MeV) and are sometimes referred to as "warm" fusion reactions. This leads, in part, to relatively high yields from these reactions.
3
Ununtrium
4
237
Np(48Ca,xn)285-x113 (x=3)
In June 2006, the Dubna-Livermore team synthesised ununtrium directly in the "warm" fusion reaction between neptunium-237 and calcium-48 nuclei. Two atoms of 282Uut were detected with a cross section of 900 fb.[10]
Chronology of isotope discovery Isotope
Year discovered
discovery reaction
278
2004
279
unknown
280
unknown
281
unknown
282
2006
237
283
2003
243
284
2003
243
Uut Uut Uut Uut Uut Uut Uut
209
Bi(70Zn,n)
[4]
[10]
Np(48Ca,3n)
[2] Am(48Ca,4n) [2] Am(48Ca,3n)
Yields of isotopes Cold fusion The table below provides cross-sections and excitation energies for cold fusion reactions producing ununtrium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
70
CN
209
Zn
279
Bi
Uut
1n
2n
3n
31 fb
Hot Fusion The table below provides cross-sections and excitation energies for hot fusion reactions producing ununtrium isotopes directly. Data in bold represents maxima derived from excitation function measurements. + represents an observed exit channel. Projectile 48
Ca
Target 237
Np
CN 285
Uut
3n 0.9 pb , 39.1 MeV
[10]
4n
5n
Ununtrium
5
Theoretical calculations on evaporation cross sections The below table contains various targets-projectile combinations for which calculations have provided estimates for cross section yields from various neutron evaporation channels. The channel with the highest expected yield is given. DNS = Di-nuclear system ; σ = cross section Target
Projectile
CN
Channel (product)
~σ
max
Model
Ref
209
Bi
70
Zn
279
1n (278113)
30 fb
DNS
[11]
237
Np
48
Ca
285
3n (282113)
0.4 pb
DNS
[12]
113 113
See also • Isotopes of ununtrium • Island of stability
External links • • • • •
WebElements.com - Uut [13] Uut and Uup Add Their Atomic Mass to Periodic Table Apsidium - Ununtrium [15] Discovery of Elements 113 and 115 [16] Superheavy elements [17]
[14]
References [1] http:/ / media. iupac. org/ publications/ pac/ 1979/ pdf/ 5102x0381. pdf [2] "Experiments on the synthesis of element 115 in the reaction 243Am(48Ca,xn)291-x115" (http:/ / www. jinr. ru/ publish/ Preprints/ 2003/ 178(E7-2003-178). pdf), Oganessian et al.., JINR Preprints, 2003. Retrieved on 2008-03-03 [3] "Experiments on the synthesis of element 115 in the reaction 243Am(48Ca,xn)291-x115" (http:/ / prola. aps. org/ abstract/ PRC/ v69/ i2/ e021601), Oganessian et al., Phys. Rev. C 69, 021601 (2004). Retrieved on 2008-03-03 [4] "Experiment on the Synthesis of Element 113 in the Reaction 209Bi(70Zn, n)278113" (http:/ / jpsj. ipap. jp/ link?JPSJ/ 73/ 2593/ ), Morita et al.., J. Phys. Soc. Jpn., 2004, 73, 10. Retrieved on 2008-03-03 [5] "RESULTS OF THE EXPERIMENT ON CHEMICAL IDENTIFICATION OF Db AS A DECAY PRODUCT OF ELEMENT 115" (http:/ / www. jinr. ru/ publish/ Preprints/ 2004/ 157(e12-2004-157). pdf), Oganessian et al., JINR preprints, 2004. Retrieved on 2008-03-03 [6] "Synthesis of elements 115 and 113 in the reaction 243Am + 48Ca" (http:/ / scitation. aip. org/ getabs/ servlet/ GetabsServlet?prog=normal& id=PRVCAN000072000003034611000001& idtype=cvips& gifs=yes), Oganessian et al., Phys. Rev. C72, 034611 (2005). Retrieved on 2008-03-03 [7] P. Roy Chowdhury, D. N. Basu and C. Samanta (2007). " α decay chains from element 113 (http:/ / scitation. aip. org/ getabs/ servlet/ GetabsServlet?prog=normal& id=PRVCAN000075000004047306000001& idtype=cvips& gifs=yes)". Phys. Rev. C 75: 047306. doi: 10.1103/PhysRevC.75.047306 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 75. 047306). . [8] " RIKEN NEWS November 2004 (http:/ / www. riken. go. jp/ engn/ r-world/ info/ release/ news/ 2004/ nov/ index. html)". . Retrieved 2008-02-09. [9] "Search for element 113" (http:/ / www. gsi. de/ informationen/ wti/ library/ scientificreport2003/ files/ 1. pdf), Hofmann et al., GSI report 2003. Retrieved on 2008-03-03 [10] "Synthesis of the isotope 282113 in the 237Np+48Ca fusion reaction" (http:/ / nrv. jinr. ru/ pdf_file/ PhysRevC_76_011601. pdf), Oganessian et al., Phys. Rev. C76, 011601(R) (2007). Retrieved on 2009-08-17 [11] http:/ / arxiv. org/ pdf/ 0707. 2588 [12] http:/ / arxiv. org/ pdf/ 0803. 1117 [13] http:/ / www. webelements. com/ webelements/ elements/ text/ Uut/ index. html
Ununtrium [14] [15] [16] [17]
http:/ / www. radiochemistry. org/ periodictable/ elements/ 115. html http:/ / www. apsidium. com/ elements/ 113. htm http:/ / www-cms. llnl. gov/ e113_115/ images. html http:/ / physicsweb. org/ articles/ world/ 17/ 7/ 7
6
Article Sources and Contributors
Article Sources and Contributors Ununtrium Source: http://en.wikipedia.org/w/index.php?oldid=308459660 Contributors: Ahoerstemeier, Alfio, AlimanRuna, Allen3, Anoop.m, Attackman7171, BlueEarth, Brett Dunbar, Bryan Derksen, Bubbha, Cacahueten, Carbuncle, Crystal whacker, Daniel bg, Darrien, David Shay, Db099221, Devleenasamanta, Dogposter, Doulos Christos, Drjezza, Edgar181, Emperorbma, Eric119, Feeeshboy, Femto, Ferengi, Flying Jazz, Fonzy, Fukumoto, Fvw, Gveret Tered, Harro5, Hashar, Hugo-cs, Icairns, Iwoelbern, Jeltz, Joanjoc, John, KJS77, Kelovy, Kingdon, Kurykh, Kwamikagami, Leapmark, Liltibs, Mav, Mellery, Merovingian, Nergaal, Netdragon, NewEnglandYankee, Nick Levine, Nozzleman, Ortolan88, Pascal666, Poohbee13666, Poolkris, Pras, Rathlan, Reddi, Remember, Reyk, Rich Farmbrough, RickK, Rickjpelleg, Rifleman 82, Roentgenium111, Romanskolduns, Rufty, Rursus, Saperaud, Schneelocke, ShaunMacPherson, Shimgray, SimonP, Siroxo, Skatebiker, Sl, Snoyes, Tagishsimon, Tetracube, Timc, Trollminator, Vicki Rosenzweig, Xoloz, Xtreambar, Zinc90, 60 anonymous edits
Image Sources, Licenses and Contributors image:Uut-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Uut-TableImage.png License: GNU Free Documentation License Contributors: Kwamikagami, Saperaud File:Speaker Icon.svg Source: http://en.wikipedia.org/w/index.php?title=File:Speaker_Icon.svg License: Public Domain Contributors: Blast, G.Hagedorn, Mobius, 2 anonymous edits
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
7
Ununquadium
1
Ununquadium ununtrium ← ununquadium → ununpentium Pb ↑ Uuq ↓ (Uhq) Periodic Table Extended Periodic Table General Name, symbol, number
ununquadium, Uuq, 114
Element category
Unknown, suspected to be a noble gas
Group, Period, Block
14, 7, p
Appearance
unknown Standard atomic weight
−1
[289] g·mol
Electron configuration
14
10
2
2
perhaps [Rn] 5f 6d 7s 7p (guess based on lead)
Electrons per shell
2, 8, 18, 32, 32, 18, 4
Phase
unknown CAS registry number
54085-16-4 Most-stable isotopes
Main article: Isotopes of ununquadium iso
NA
half-life
DM
DE (MeV)
DP
syn
2.6 s
α
9.82,9.48
285
289b
syn
1.1 m
α
9.67
285b
288
syn
0.8 s
α
9.94
284
287
syn
0.48 s
α
10.02
283
287b
syn
5.5 s
α
10.29
283b
286
syn
0.13 s
40% α
10.19
282
289
Uuq Uuq ?
Uuq Uuq Uuq ??
Uuq
Uub Uub ?
Uub Uub Uub ??
Uub
60% SF References
Ununquadium (pronounced /ˌjuːnənˈkwɒdiəm/;[1] officially, the two initial u's are to be [2] pronounced English pronunciation: /uː/ ( listen) ) is the temporary name of a radioactive
Ununquadium
2
chemical element in the periodic table that has the temporary symbol Uuq and has the atomic number 114. About 80 decays of atoms of ununquadium have been observed to date, 50 directly and 30 from the decay of the heavier elements ununhexium and ununoctium. All decays have been assigned to the four neighbouring isotopes with mass numbers 286-289. The longest-lived isotope currently known is 289114 with a half-life of ~2.6 s, although there is evidence for an isomer, 289b114, which has a half-life of ~66 s, which would represent one of the longest-lived nuclei in the superheavy element region. Recent chemistry experiments have strongly indicated that element 114 possesses non-'eka'-lead properties and appears to behave as the first superheavy element that portrays noble-gas-like properties due to relativistic effects.[3]
De facto discovery In December 1998, scientists at Dubna (Joint Institute for Nuclear Research) in Russia bombarded a Pu-244 target with Ca-48 ions. A single atom of element 114, decaying by 9.67 MeV alpha-emission with a half-life of 30 s, was produced and assigned to 289114. This observation was subsequently published in January 1999.[4] However, the decay chain observed has not been repeated and the exact identity of this activity is unknown although it is possible that it is due to a meta-stable isomer, namely 289m114. In March 1999, the same team replaced the Pu-244 target with a Pu-242 one in order to produce other isotopes. This time two atoms of element 114 were produced, decaying by 10.29 MeV alpha-emission with a half-life of 5.5 s. They were assigned as 287114.[5] Once again, this activity has not been seen again and it is unclear what nucleus was produced. It is possible that it was a meta-stable isomer, namely 287m114. The now-confirmed discovery of element 114 was made in June 1999 when the Dubna team repeated the Pu-244 reaction. This time, two atoms of element 114 were produced decaying by emission of 9.82 MeV alpha particles with a half life of 2.6 s.[6] This activity was initially assigned to 288114 in error, due to the confusion regarding the above observations. Further work in Dec 2002 has allowed a positive reassignment to 289 114.[7] 24494Pu
+
4820Ca
→
292114Uuq
*
→
289114Uuq
+3
10n
Theoretical estimation of the alpha decay half lives of the isotopes of the element 114 supports the experimental data.[8] [9] The fission-survived isotope 298114 is predicted to have alpha decay half life around 17 days.[10] [11] In May 2009, the JWP published a report on the discovery of element 112 ununbium in which they acknowledged the discovery of the isotope 283112.[12] This therefore implies the de facto discovery of element 114, from the acknowledgment of the data for the synthesis of 287 114 and 291116 (see below), relating to 283112, although this may not be determined as the first synthesis of the element. An impending report by the JWP will discuss these issues. The discovery of element 114, as 2009 at the GSI (see section 6).
288
114 and
289
114, was successfully confirmed in July
Ununquadium
Naming Current names Ununquadium (Uuq) is a temporary IUPAC systematic element name. Research scientists usually refer to the element simply as element 114.
Proposed names by claimants According to IUPAC recommendations, the discoverer(s)of a new element has the right to suggest a name.[13] No naming suggestions have yet been given by the (claimant) discoverers.
Extrapolated chemical properties of eka-lead Oxidation states Element 114 is projected to be the second member of the 7p series of non-metals and the heaviest member of group 14 (IVA) in the Periodic Table, below lead. Each of the members of this group show the group oxidation state of +IV and the latter members have an increasing +II chemistry due to the onset of the inert pair effect. Tin represents the point at which the stability of the +II and +IV states are similar. Lead, the heaviest member, portrays a switch from the +IV state to the +II state. Element 114 should therefore follow this trend and a possess an oxidising +IV state and a stable +II state.
Chemistry Element 114 should portray eka-lead chemical properties and should therefore form a monoxide, UuqO, and dihalides, UuqF2, UuqCl2, UuqBr2, and UuqI2. If the +IV state is accessible, it is likely that it is only possible in the oxide, UuqO2, and fluoride, UuqF4. It may also show a mixed oxide, Uuq3O4, analogous to Pb3O4. Some studies also suggest that the chemical behaviour of element 114 might in fact be closer to that of the noble gas radon, than to that of lead.[3]
Experimental chemistry Atomic gas phase Two experiments were performed in April-May 2007 in a joint FLNR-PSI collaboration aiming to study the chemistry of element 112. The first experiment involved the reaction 242 Pu(48Ca,3n)287114 and the second the reaction 244Pu(48Ca,4n)288114. The adsorption properties of the resultant atoms on a gold surface were compared with those of radon. The first experiment allowed detection of 3 atoms of 283112 (see ununbium) but also seemingly detected 1 atom of 287114. This result was a surprise given the transport time of the product atoms is ~2 s, so element 114 atoms should decay before adsorption. In the second reaction, 2 atoms of 288114 and possibly 1 atom of 289114 were detected. Two of the three atoms portrayed adsorption characteristics associated with a volatile, noble-gas-like element, which has been suggested but is not predicted by more recent calculations. These experiments did however provide independent confirmation for the discovery of elements 112, 114, and 116 via comparison with published decay data. Further experiments were
3
Ununquadium performed in 2008 to confirm this important result and a single atom of 289114 was detected which gave data in agreement with previous data in support of element 114 having a noble-gas-like interaction with gold.[14]
History of synthesis of isotopes by cold fusion This section deals with the synthesis of nuclei of ununquadium by so-called "cold" fusion reactions. These are processes which create compound nuclei at low excitation energy (~10-20 MeV, hence "cold"), leading to a higher probability of survival from fission. The excited nucleus then decays to the ground state via the emission of one or two neutrons only. 208
Pb(76Ge,xn)284−x114
The first attempt to synthesise element 114 in cold fusion reactions was performed at GANIL, France in 2003. No atoms were detected providing a yield limit of 1.2 pb.
History of synthesis of isotopes by hot fusion This section deals with the synthesis of nuclei of ununquadium by so-called "hot" fusion reactions. These are processes which create compound nuclei at high excitation energy (~40-50 MeV, hence "hot"), leading to a reduced probability of survival from fission. The excited nucleus then decays to the ground state via the emission of 3-5 neutrons. Fusion reactions utilizing 48Ca nuclei usually produce compound nuclei with intermediate excitation energies (~30-35 MeV) and are sometimes referred to as "warm" fusion reactions. This leads, in part, to relatively high yields from these reactions. 244
Pu(48Ca,xn)292−x114 (x=3,4,5)
The first experiments on the synthesis of element 114 were performed by the team in Dubna in November 1998. They were able to detect a single, long decay chain, assigned to 289 114.[4] The reaction was repeated in 1999 and a further 2 atoms of element 114 were 288 [6] detected. The products were assigned to 114. The team further studied the reaction in 2002. During the measurement of the 3n, 4n, and 5n neutron evaporation excitation functions they were able to detect 3 atoms of 289114, 12 atoms of the new isotope 288114, and 1 atom of the new isotope 287114. Based on these results, the first atom to be detected was tentatively reassigned to 290114 or 289m114, whilst the two subsequent atoms were reassigned to 289114 and therefore belong to the unofficial discovery experiment.[7] In an attempt to study the chemistry of element 112 as the isotope 285112, this reaction was repeated in April 2007. Surprisingly, a PSI-FLNR directly detected 2 atoms of 288114 forming the basis for the first chemical studies of element 114. In June 2008, the experiment was repeated in order to further assess the chemistry of the element using the 289 114 isotope. A single atom was detected seeming to confirm the noble-gas-like properties of the element. During May-July 2009, the team at GSI studied this reaction for the first time, as a first step towards the synthesis of element 117.[15] . The team were able to confirm the synthesis and decay data for 288114 and 289114.[16]
4
Ununquadium
5
242
Pu(48Ca,xn)290−x114 (x=2,3,4)
The team at Dubna first studied this reaction in March-April 1999 and detected two atoms of element 114, assigned to 287114.[5] The reaction was repeated in September 2003 in order to attempt to confirm the decay data for 287114 and 283112 since conflicting data for 283 112 had been collected (see ununbium). The Russian scientists were able to measure decay data for 288114,287114 and the new isotope 286114 from the measurement of the 2n, 3n, and 4n excitation functions. [17] [18] In April 2006, a PSI-FLNR collaboration used the reaction to determine the first chemical properties of element 112 by producing 283112 as an overshoot product. In a confirmatory experiment in April 2007, the team were able to detect 287114 directly and therefore measure some initial data on the atomic chemical properties of element 114.
Synthesis of isotopes as decay products The isotopes of ununquadium have also been observed in the decay of elements 116 and 118 (see ununoctium for decay chain). Evaporation residue
Observed Uuq isotope
293
289
[19] [18]
292
288
[18]
291
287
[7]
286
[20]
116
114
116
114
116
294
118,
114
290
116
114
Chronology of isotope discovery Isotope
Year discovered
Discovery reaction
2002
249
2002
244
287b
1999
242
288
2002
244
1999
244
1998
244
286
Uuq
287a
Uuq Uuq ??
Uuq
289a
Uuq
289b
Uuq ?
Cf(48Ca,3n) Pu(48Ca,5n) Pu(48Ca,3n) Pu(48Ca,4n) Pu(48Ca,3n) Pu(48Ca,3n)
[20]
Ununquadium
6
Yields of isotopes The tables below provide cross-sections and excitation energies for fusion reactions producing ununquadium isotopes directly. Data in bold represent maxima derived from excitation function measurements. + represents an observed exit channel.
Cold fusion Projectile
Target
76
208
Ge
CN
2n
3n
< 1.2 pb
284
Pb
1n
Uuq
Hot fusion Projectile
Target
CN
2n
3n
4n
5n
48
242
290
0.5 pb, 32.5 MeV 3.6 pb, 40.0 MeV
4.5 pb, 40.0 MeV
<1.4 pb , 45.0 MeV
48
244
292
1.7 pb, 40.0 MeV
5.3 pb, 40.0 MeV
1.1 pb, 52.0 MeV
Ca
Ca
Pu
Pu
Uuq
Uuq
Isomerism in ununquadium isotopes 289
114
In the first claimed synthesis of element 114, an isotope assigned as 289114 decayed by emitting a 9.71 MeV alpha particle with a lifetime of 30 seconds. This activity was not observed in repetitions of the direct synthesis of this isotope. However, in a single case from the synthesis of 293116, a decay chain was measured starting with the emission of a 9.63 MeV alpha particle with a lifetime of 2.7 minutes. All subsequent decays were very similar to that observed from 289114, presuming that the parent decay was missed. This strongly suggests that the activity should be assigned to an isomeric level. The absence of the activity in recent experiments indicates that the yield of the isomer is ~20% compared to the supposed ground state and that the observation in the first experiment was a fortunate (or not as the case history indicates). Further research is required to resolve these issues. 287
114
In a manner similar to those for 289114, first experiments with a 242Pu target identified an isotope 287114 decaying by emission of a 10.29 MeV alpha particle with a lifetime of 5.5 seconds. The daughter spontaneously fissioned with a lifetime in accord with the previous synthesis of 283112. Both these acitivities have not been observed since (see ununbium). However, the correlation suggests that the results are not random and are possible due to the formation of isomers whose yield is obviously dependent on production methods. Further research is required to unravel these discrepancies.
Ununquadium
7
Retracted isotopes 285
114
In the claimed synthesis of 293118 in 1999, the isotope 285114 was identified as decaying by 11.35 MeV alpha emission with a half-life of 0.58 ms. The claim was retracted in 2001 and hence this ununquadium isotope is currently unknown or unconfirmed.
Fission of compound nuclei with Z=114 Several experiments have been performed between 2000-2004 at the Flerov Laboratory of Nuclear Reactions in Dubna studying the fission characteristics of the compound nucleus 292 114. The nuclear reaction used is 244Pu+48Ca. The results have revealed how nuclei such as this fission predominantly by expelling closed shell nuclei such as 132Sn (Z=50, N=82). It was also found that the yield for the fusion-fission pathway was similar between 48 Ca and 58Fe projectiles, indicating a possible future use of 58Fe projectiles in superheavy element formation.[21]
Target-Projectile Combinations leading to Z=114 compound nuclei The below table contains various combinations of targets and projectiles which could be used to form compound nuclei with Z=114. A indicates a successful reaction, a indicates a failure to date and a indicates a reaction yet to be attempted.
Target-projectile combinations leading to Z=116 nuclei Target
Projectile
208
Pb
76
232
Th
54
238
Ge Cr
CN 284
114
286
114
U
50
Ti
288
244
Pu
48
Ca
292
242
48
290
239
Pu
48
287
248
Cm
40
Pu
249
Cf
Ca Ca Ar
36
S
114 114 114 114
288
114
285
114
Attempted?
Ununquadium
8
Theoretical calculations on evaporation cross sections The below table contains various targets-projectile combinations for which calculations have provided estimates for cross section yields from various neutron evaporation channels. The channel with the highest expected yield is given. DNS = Dinuclear system ; σ = cross section Target
Projectile
CN
Channel (product)
~σ
max
Model
Ref
208
76
284
1n (283114)
60 fb
DNS
[22]
208
73
281
1n (280114)
0.2 pb
DNS
[22]
Pb Pb
Ge Ge
114 114
U
50
Ti
288
2n (286114)
60 fb
DNS
[23]
244
Pu
48
Ca
292
4n (288114)
5 pb
DNS
[23]
242
48
290
4n (286114)
0.4 pb
DNS
[23]
238
Pu
Ca
114 114 114
In search for the island of stability:
298
114
According to macroscopic-microscopic (MM) theory, Z=114 is the next spherical magic number. This means that such nuclei are spherical in their ground state and should have high, wide fission barriers to deformation and hence long SF partial half-lives. In the region of Z=114, MM theory indicates that N=184 is the next spherical neutron magic number and puts forward the nucleus 298114 as a strong candidate for the next spherical doubly magic nucleus, after 208Pb (Z=82, N=126). 298114 is taken to be at the centre of a hypothetical ‘island of stability’. However, other calculations using relativistic mean field (RMF) theory propose Z=120, 122, and 126 as alternative proton magic numbers depending upon the chosen set of parameters. It is possible that rather than a peak at a specific proton shell, there exists a plateau of proton shell effects from Z=114–126. It should be noted that calculations suggest that the minimum of the shell-correction energy and hence the highest fission barrier exists for 297115, caused by pairing effects. Due to the expected high fission barriers, any nucleus within this island of stability will exclusively decay by alpha-particle emission and as such the nucleus with the longest half-life is predicted to be 298114. The expected half-life is unlikely to reach values higher than about 10 minutes, unless the N=184 neutron shell proves to be more stabilising than predicted, for which there exists some evidence. In addition, 297114 may have an even-longer half-life due to the effect of the odd neutron, creating transitions between similar Nilsson levels with lower Qalpha values. In either case, an island of stability does not represent nuclei with the longest half-lives but those which are significantly stabilized against fission by closed-shell effects.
Evidence for Z=114 closed proton shell Whilst evidence for closed neutron shells can be deemed directly from the systematic variation of Qalpha values for ground-state to ground-state transitions, evidence for closed proton shells comes from (partial) spontaneous fission half-lives. Such data can sometimes be difficult to extract due to low production rates and weak SF branching. In the case of Z=114, evidence for the effect of this proposed closed shell comes from the comparison
Ununquadium
9
between the nuclei pairings 282112 (TSF1/2 = 0.8 ms) and 286114 (TSF1/2 = 130 ms), and 284 112 (TSF = 97 ms) and 288114 (TSF >800 ms). Further evidence would come from the measurement of partial SF half-lives of nuclei with Z>114, such as 290116 and 292118 (both N=174 isotones). The extraction of Z=114 effects is complicated by the presence of a dominating N=184 effect in this region.
Difficulty of synthesis of 298114 The direct synthesis of the nucleus 298114 by a fusion-evaporation pathway is impossible since no known combination of target and projectile can provide 184 neutrons in the compound nucleus. It has been suggested that such a neutron-rich isotope can be formed by the quasifission (partial fusion followed by fission) of a massive nucleus. Such nuclei tend to fission with the formation of isotopes close to the closed shells Z=20/N=20 (40Ca), Z=50/N=82 (132Sn) or Z=82/N=126 (208Pb/209Bi). If Z=114 does represent a closed shell, then the hypothetical reaction below may represent a method of synthesis: 20480Hg
+
13654Xe
→
298114Uuq
+
4020Ca
+2
10n
298
It is also possible that 114 can be synthesized by the alpha decay of a massive nucleus. Such a method would depend highly on the SF stability of such nuclei, since the alpha half-lives are expected to be very short. The yields for such reactions will also most likely be extremely small. One such reaction is: 24494Pu(9640Zr,
2n) →
338134Utq
→→
298114Uuq
+ 10
42He
Currently running experiments In February 2009, it was indicated that the team at LBNL may have detected an atom of element 114 using the previously-studied reaction 242Pu + 48Ca.[24] Confirmation is not yet available. In addition, in April 2009, the PSI-FLNR collaboration carried out another study of the chemistry of element 114. Results are not yet available.
Future experiments The team at RIKEN have indicated plans to study the cold fusion reaction: 20882Pb
+
7632Ge
→
284114Uuq
*
→?
The TASCA collaboration based at the GSI will perform their first chemistry experiments on E114 starting in August 2009, following their successful production of the element in April 2009. The FLNR have future plans to study light isotopes of element 114, formed in the reaction between 239Pu and 48Ca.
Ununquadium
10
See also • • • •
Island of stability: Ununquadium–Unbinilium–Unbihexium Lead Periodic table (extended) Isotopes of ununquadium
External links • WebElements.com: Uuq [25] • First postcard from the island of nuclear stability • Second postcard from the island of stability [27]
[26]
References [1] ununquadium (http:/ / reference. aol. com/ columbia/ _a/ ununquadium/ 20051207161909990014) [2] http:/ / media. iupac. org/ publications/ pac/ 1979/ pdf/ 5102x0381. pdf [3] Gas Phase Chemistry of Superheavy Elements (http:/ / lch. web. psi. ch/ files/ lectures/ TexasA& M/ TexasA& M. pdf), lecture by Heinz W. Gäggeler, Nov. 2007. Last accessed on Dec. 12, 2008. [4] "Synthesis of Superheavy Nuclei in the 48Ca + 244Pu Reaction" (http:/ / prola. aps. org/ abstract/ PRL/ v83/ i16/ p3154_1), Oganessian et al.., Phys. Rev. Lett. 83, 3154 - 3157 (1999).Retrieved on 2008-03-03 [5] "Synthesis of nuclei of the superheavy element 114 in reactions induced by 48Ca" (http:/ / www. nature. com/ nature/ journal/ v400/ n6741/ abs/ 400242a0. html), Oganessian et al., Nature 400, 242-245 (15 July 1999). Retrieved on 2008-03-03 [6] "Synthesis of superheavy nuclei in the 48Ca+244Pu reaction: 288114" (http:/ / prola. aps. org/ abstract/ PRC/ v62/ i4/ e041604), Oganessian et al.., Phys. Rev. C 62, 041604 (2000). Retrieved on 2008-03-03 [7] "Measurements of cross sections for the fusion-evaporation reactions 244Pu(48Ca,xn)292−x114 and 245 Cm(48Ca,xn)293−x116" (http:/ / prola. aps. org/ abstract/ PRC/ v69/ i5/ e054607), Oganessian et al., Phys. Rev. C 69, 054607 (2004). Retrieved on 2008-03-03 [8] P. Roy Chowdhury, C. Samanta, and D. N. Basu (26 January 2006). " α decay half-lives of new superheavy elements (http:/ / scitation. aip. org/ getabs/ servlet/ GetabsServlet?prog=normal& id=PRVCAN000073000001014612000001& idtype=cvips& gifs=yes)". Phys. Rev. C 73: 014612. doi: 10.1103/PhysRevC.73.014612 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 73. 014612). . [9] C. Samanta, P. Roy Chowdhury and D.N. Basu (2007). " Predictions of alpha decay half lives of heavy and superheavy elements (http:/ / www. sciencedirect. com/ science?_ob=ArticleURL& _udi=B6TVB-4NF4F0Y-2& _user=2806701& _rdoc=1& _fmt=& _orig=search& _sort=d& view=c& _acct=C000058844& _version=1& _urlVersion=0& _userid=2806701& md5=3f680654b5659191d67f31681a4cfc83)". Nucl. Phys. A 789: 142–154. doi: 10.1016/j.nuclphysa.2007.04.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2007. 04. 001). . [10] P. Roy Chowdhury, C. Samanta, and D. N. Basu (2008). " Search for long lived heaviest nuclei beyond the valley of stability (http:/ / scitation. aip. org/ getabs/ servlet/ GetabsServlet?prog=normal& id=PRVCAN000077000004044603000001& idtype=cvips& gifs=yes)". Phys. Rev. C 77: 044603. doi: 10.1103/PhysRevC.77.044603 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 77. 044603). . [11] P. Roy Chowdhury, C. Samanta, and D. N. Basu (2008). " Nuclear half-lives for α-radioactivity of elements with 100 ≤ Z ≤ 130 (http:/ / www. sciencedirect. com/ science?_ob=ArticleURL& _udi=B6WBB-4S26JRX-1& _user=2806701& _coverDate=03/ 14/ 2008& _alid=740505626& _rdoc=6& _fmt=high& _orig=search& _cdi=6706& _sort=d& _docanchor=& view=c& _ct=211& _acct=C000058844& _version=1& _urlVersion=0& _userid=2806701& md5=dc85a3a8a2ac1faa38c3804f16f86c13)". At. Data & Nucl. Data Tables 94: 781–806. doi: 10.1016/j.adt.2008.01.003 (http:/ / dx. doi. org/ 10. 1016/ j. adt. 2008. 01. 003). . [12] R.C.Barber; H.W.Gaeggeler;P.J.Karol;H. Nakahara; E.Verdaci; E. Vogt (2009). " Discovery of the element with atomic number 112 (http:/ / media. iupac. org/ publications/ pac/ asap/ pdf/ PAC-REP-08-03-05. pdf)" (IUPAC Technical Report). Pure Appl. Chem.. doi: 10.1351/PAC-REP-08-03-05 (http:/ / dx. doi. org/ 10. 1351/ PAC-REP-08-03-05). . [13] http:/ / media. iupac. org/ publications/ pac/ 2002/ pdf/ 7405x0787. pdf [14] http:/ / www1. jinr. ru/ Reports/ 2008/ english/ 06_flnr_e. pdf [15] https:/ / www. gsi. de/ forschung/ beamtime/ Schedule/ 2009/ Month/ Mo05. pdf [16] http:/ / www-w2k. gsi. de/ tasca/ News_Dates/ news_dates2. htm
Ununquadium [17] "Measurements of cross sections and decay properties of the isotopes of elements 112, 114, and 116 produced in the fusion reactions 233,238U, 242Pu, and 248Cm+48Ca" (http:/ / prola. aps. org/ abstract/ PRC/ v70/ i6/ e064609), Oganessian et al., Phys. Rev. C 70, 064609 (2004). Retrieved on 2008-03-03 [18] "Measurements of cross sections and decay properties of the isotopes of elements 112, 114, and 116 produced in the fusion reactions 233,238U , 242Pu , and 248Cm+48Ca" (http:/ / www. jinr. ru/ publish/ Preprints/ 2004/ 160(E7-2004-160). pdf), Oganessian et al., JINR preprints, 2004. Retrieved on 2008-03-03 [19] [20] [21] [22] [23] [24] [25] [26] [27]
see ununhexium see ununoctium see Flerov lab annual reports 2000-2006 inclusive http:/ / www1. jinr. ru/ Reports/ Reports_eng_arh. html http:/ / arxiv. org/ pdf/ 0707. 2588 http:/ / arxiv. org/ pdf/ 0803. 1117 http:/ / www. chemicalforums. com/ index. php?topic=30758. 0 http:/ / webelements. com/ webelements/ elements/ text/ Uuq/ index. html http:/ / cerncourier. com/ main/ article/ 39/ 7/ 18 http:/ / www. cerncourier. com/ main/ article/ 41/ 8/ 17
11
Article Sources and Contributors
Article Sources and Contributors Ununquadium Source: http://en.wikipedia.org/w/index.php?oldid=308005746 Contributors: Adashiel, Ahoerstemeier, AkvoD3, Alfio, AlimanRuna, Angstengel, Archfalhwyl, AxelBoldt, Bhangranuch, BlueEarth, Bobblewik, Bryan Derksen, Cacahueten, Carbuncle, ChemNerd, Cosmium, Dajwilkinson, Dale101usa, Dante Alighieri, Darrien, David Shay, Devleenasamanta, Drjezza, Edgar181, El C, Emperorbma, Eric119, Erkcan, Femto, Flying Jazz, Fonzy, Fosnez, Fvw, GDonato, Green Bush Draw, Gveret Tered, Hashar, Headbomb, Hugo-cs, Icairns, Icarus3, Jeffreymcmanus, Joanjoc, John, Jose Ramos, Keenan Pepper, Kelovy, Kingdon, Kurt Jansson, Kurykh, Kwamikagami, MARKELLOS, Manoj C Menon, Markhurd, Mav, Nergaal, Nihiltres, Noisy, Nozzleman, Orelstrigo, Ortolan88, Oscarcwk, Ossipewsk, Otto ter Haar, Parsa, Pascal666, Pewwer42, Pion, Polonium, Poolkris, Potatoswatter, Pt, Qoou.Anonimu, RTC, Rathlan, Remember, Reyk, Rich Farmbrough, RickK, Rickjpelleg, Rifleman 82, Rjwilmsi, Roberta F., Roentgenium111, Romanm, Romanskolduns, Rumping, Rursus, Saperaud, ShaunMacPherson, SimonP, Siroxo, Skatebiker, Skomorokh, Sl, Stirling Newberry, TUF-KAT, Tagishsimon, Tetracube, Theseeker4, Timc, Timwi, Tonyrex, Trollminator, Typhoonchaser, Vicki Rosenzweig, Vuo, Wafulz, Walkerma, Warut, Zkamran53, 118 anonymous edits
Image Sources, Licenses and Contributors image:Uuq-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Uuq-TableImage.png License: GNU Free Documentation License Contributors: Kwamikagami, Saperaud File:Speaker Icon.svg Source: http://en.wikipedia.org/w/index.php?title=File:Speaker_Icon.svg License: Public Domain Contributors: Blast, G.Hagedorn, Mobius, 2 anonymous edits Image:Yes check.svg Source: http://en.wikipedia.org/w/index.php?title=File:Yes_check.svg License: GNU Free Documentation License Contributors: User:Gmaxwell, User:WarX Image:X mark.svg Source: http://en.wikipedia.org/w/index.php?title=File:X_mark.svg License: GNU Free Documentation License Contributors: Abnormaal, Gmaxwell, Kilom691, MGA73, Mardetanha, Penubag, Pseudomoi, WikipediaMaster, 1 anonymous edits Image:Black x.svg Source: http://en.wikipedia.org/w/index.php?title=File:Black_x.svg License: unknown Contributors: User:Howcheng
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12
Ununpentium
1
Ununpentium ununquadium ← ununpentium → ununhexium Bi ↑ Uup ↓ (Uhp) Periodic Table Extended Periodic Table General Name, symbol, number
ununpentium, Uup, 115
Element category
presumably poor metals
Group, Period, Block
15, 7, p
Standard atomic weight
−1
[288] g·mol
Electron configuration
14
10
2
3
perhaps [Rn] 5f 6d 7s 7p (guess based on bismuth)
Electrons per shell
2, 8, 18, 32, 32, 18, 5
CAS registry number
54085-64-2 Most-stable isotopes
Main article: Isotopes of ununpentium iso
NA
half-life
288
syn
87.5 ms
287
syn
32 ms
Uup Uup
DM α α
DE (MeV)
DP
10.46
284
10.59
283
Uut Uut
References
Ununpentium (pronounced /ˌjuːnənˈpɛntiəm/; officially, the two initial u's are to be [1] pronounced English pronunciation: /uː/ ( listen) ) is the temporary name of a synthetic superheavy element in the periodic table that has the temporary symbol Uup and has the atomic number 115. Only about 30 atoms of ununpentium have been synthesised to date and just 4 direct decays of the parent element have been detected. Two isotopes are currently known, Uup-287 and Uup-288. Element 115 also falls in the center of the theoretical island of stability. The most stable isotope of ununpentium is predicted to be Uup-299, containing the theorized "magic number" of 184 neutrons. The most neutron rich isotope to date is Uup-288, which contains 173 neutrons.
Ununpentium
Discovery profile On February 2, 2004, synthesis of ununpentium was reported in Physical Review C by a team composed of Russian scientists at the Joint Institute for Nuclear Research in Dubna, and American scientists at the Lawrence Livermore National Laboratory.[2] [3] The team reported that they bombarded americium-243 with calcium-48 ions to produce four atoms of ununpentium. These atoms, they report, decayed by emission of alpha-particles to ununtrium in approximately 100 milliseconds.
The Dubna-Livermore collaboration has strengthened their claim for the discovery of ununpentium by conducting chemical experiments on the decay daughter 268Db. In experiments in June 2004 and December 2005, the Dubnium isotope was successfully identified by milking the Db fraction and measuring any SF activities.[4] [5] Both the half-life and decay mode were confirmed for the proposed 268Db which lends support to the assignment of Z=115 to the parent nuclei. Theoretical calculations using a quantum-tunneling model support the experimental alpha-decay half-lives.[6]
Official claim of discovery of element 115 Sergei Dmitriev from the Flerov laboratory of nuclear reactions (FLNR) in Dubna, Russia, has formally put forward their claim of discovery of element 115 to the Joint Working Party (JWP) from IUPAC and IUPAP. The JWP are expected to publish their opinions on such claims in the near future.[7]
Naming Current names Ununpentium is historically known as eka-bismuth. Ununpentium is a temporary IUPAC systematic element name. Research scientists usually refer to the element simply as element 115.
Extrapolated chemical properties of eka-bismuth Oxidation states Element 115 is projected to be the third member of the 7p series of non-metals and the heaviest member of group 15 (VA) in the Periodic Table, below bismuth. In this group, each member is known to portray the group oxidation state of +V but with differing stability. For nitrogen, the +V state is very difficult to achieve due to the lack of low-lying d-orbitals and the inability of the small nitrogen atom to accommodate five ligands. The +V state is well represented for phosphorus, arsenic, and antimony. However, for bismuth it is rare due to the reluctance of the 6s2 electron to participate in bonding. This effect is known as the "inert pair effect" and is commonly linked to relativistic stabilisation of the 6s-orbitals. It is expected that element 115 will continue this trend and portray only +III and +I oxidation states. Nitrogen(I) and bismuth(I) are known but rare and Uup(I) is likely to show some unique properties.[8]
2
Ununpentium
3
Chemistry It is expected that the chemistry of ununpentium will be related to its lighter homologue bismuth. In this regard it is expected to undergo oxidation only as far as the trioxide Uup2O3. Oxidation with the more reactive halogens should form the trihalides, such as UupF3 and UupCl3. The less-oxidising, heavier halogens, may well only be able to promote the formation of the monohalides, UupBr and UupI.
History of synthesis of isotopes by hot fusion This section deals with the synthesis of nuclei of ununpentium by so-called "hot" fusion reactions. These are processes which create compound nuclei at high excitation energy (~40-50 MeV, hence "hot"), leading to a reduced probability of survival from fission. The excited nucleus then decays to the ground state via the emission of 3-5 neutrons. Fusion reactions utilizing 48Ca nuclei usually produce compound nuclei with intermediate excitation energies (~30-35 MeV) and are sometimes referred to as "warm" fusion reactions. This leads, in part, to relatively high yields from these reactions. 238
U(51V,xn)289−xUup
There are strong indications that this reaction was performed in late 2004 as part of a uranium(IV) fluoride target test at the GSI. No reports have been published suggesting that no products atoms were detected, as anticipated by the team.[9] 243
Am(48Ca,xn)291−xUup (x=3,4)
This reaction was first performed by the team in Dubna in July-August 2003. In two separate runs they were able to detect 3 atoms of 288Uup and a single atom of 287Uup. The reaction was studied further in June 2004 in an attempt to isolate the descendant 268Db from the 288Uup decay chain. After chemical separation of a +4/+5 fraction, 15 SF decays were measured with a lifetime consistent with 268Db. In order to prove that the decays were from dubnium-268, the team repeated the reaction in August 2005 and separated the +4 and +5 fractions and further separated the +5 fractions into tantalum-like and niobium-like ones. Five SF activities were observed, all occurring in the +5 fractions and none in the tantalum-like fractions, proving that the product was indeed isotopes of dubnium.
Chronology of isotope discovery Isotope
Year discovered
Discovery reaction
287
2003
243
288
2003
243
Uup Uup
Am(48Ca,4n) Am(48Ca,3n)
Ununpentium
4
Yields of isotopes Hot fusion The table below provides cross-sections and excitation energies for hot fusion reactions producing ununpentium isotopes directly. Data in bold represent maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
48
243
Ca
CN
2n
4n
3.7 pb, 39.0 MeV
291
Am
3n
Uup
5n
0.9 pb, 44.4 MeV
Target-Projectile Combinations leading to Z=115 compound nuclei The below table contains various combinations of targets and projectiles which could be used to form compound nuclei with Z=115. A indicates a successful reaction, a indicates a failure to date and a indicates a reaction yet to be attempted.
Target-projectile combinations leading to Z=115 nuclei Target
Projectile
CN
208
Pb
75
As
283
232
Th
55
Mn
287
238
115 115
U
51
V
289
Np
50
Ti
287
244
Pu
45
Sc
289
243
Am
48
Ca
291
241
48
289
237
Am
115 115 115 115
Ca
248
115
41
Cm
249
V
289
Ar
289
115
40
Bk
249
115
37
Cf
Attempted?
286
Cl
115
Theoretical calculations on evaporation cross sections The below table contains various targets-projectile combinations for which calculations have provided estimates for cross section yields from various neutron evaporation channels. The channel with the highest expected yield is given. DNS = Di-nuclear system ; σ = cross section Target
Projectile
CN
Channel (product)
~σ
max
Model
Ref
243
48
291
3n (288115)
1 pb
DNS
[10]
242
48
290
3n (287115)
2.5 pb
DNS
[10]
Am Am
Ca Ca
115 115
Ununpentium
5
Future experiments The team at RIKEN are planning to study the reaction
As a primary next-goal for the Dubna team, they are planning to examine to products of the 243 Am + 48Ca using mass spectrometry in their state-of-the-art MASHA machine. They will attempt to isolate the dubnium products, convert them chemically into a volatile compound, most likely 268DbCl5, and measure the mass directly. The FLNR also have future plans to study light isotopes of element 115 using the reaction 241 Am + 48Ca.[11]
See also • Island of stability • Elerium-115
External links • • • •
Uut and Uup Add Their Atomic Mass to Periodic Table Apsidium: Ununpentium [13] Superheavy elements [14] History and etymology [15]
[12]
References [1] J. Chatt (IUPAC). " Recommendations for the naming of elements of atomic number greater than 100 (http:/ / media. iupac. org/ publications/ pac/ 1979/ pdf/ 5102x0381. pdf)". Pure & Appi. Chem. 51: 381. . [2] Oganessian, Yu. Ts. (2004). "Experiments on the synthesis of element 115 in the reaction 243 Am(48Ca,xn)291?x115". Physical Review C 69: 021601. doi: 10.1103/PhysRevC.69.021601 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 69. 021601). [3] Oganessian et al. (2003). "[http://www.jinr.ru/publish/Preprints/2003/178(E7-2003-178).pdf Experiments on the synthesis of element 115 in the reaction 243Am(48Ca,xn)291−x115"]]". JINR preprints. http:/ / www. jinr. ru/ publish/ Preprints/ 2003/ 178(E7-2003-178). pdf. [4] Oganessian et al. (2004). " Results of the experiment on chemical identification of db as a decay product of element 115 (http:/ / www. jinr. ru/ publish/ Preprints/ 2004/ 157(e12-2004-157). pdf)". JINR preprints. . [5] Oganessian, Yu. Ts. (2005). "Synthesis of elements 115 and 113 in the reaction ^{243}Am+^{48}Ca". Physical Review C 72: 034611. doi: 10.1103/PhysRevC.72.034611 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 72. 034611). [6] C. Samanta, P. Roy Chowdhury and D.N. Basu (2007). "Predictions of alpha decay half lives of heavy and superheavy elements". Nucl. Phys. A 789: 142–154. doi: 10.1016/j.nuclphysa.2007.04.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2007. 04. 001). [7] " Project: Priority claims for the discovery of elements with atomic number greater than 111 (http:/ / www. iupac. org/ web/ ins/ 2006-046-1-200)". IUPAC. . Retrieved 2009-07-07. [8] Keller, O. L., Jr.; C. W. Nestor, Jr. (1974). "Predicted properties of the superheavy elements. III. Element 115, Eka-bismuth". Journal of Physical Chemistry 78: 1945. doi: 10.1021/j100612a015 (http:/ / dx. doi. org/ 10. 1021/ j100612a015). [9] " List of experiments 2000-2006 (http:/ / opal. dnp. fmph. uniba. sk/ ~beer/ experiments. php)". . [10] http:/ / arxiv. org/ pdf/ 0803. 1117 [11] " Study of heavy and superheavy nuclei (see experiment 1.5) (http:/ / flerovlab. jinr. ru/ flnr/ education_list. html)". . [12] http:/ / radiochemistry. org/ periodictable/ elements/ 115. html [13] http:/ / apsidium. com/ elements/ 115. htm [14] http:/ / physicsweb. org/ articles/ world/ 17/ 7/ 7 [15] http:/ / vanderkrogt. net/ elements/ elem/ uup. html
Article Sources and Contributors
Article Sources and Contributors Ununpentium Source: http://en.wikipedia.org/w/index.php?oldid=307818616 Contributors: AKismet, Adamfinmo, Adashiel, Ahoerstemeier, Alex S, Alfio, AlimanRuna, Andareed, Animus999, Ann Stouter, Anomie, Are you ready for IPv6?, Avono, BlueEarth, Bobblewik, Bornhj, Brett Dunbar, Bryan Derksen, Cacahueten, Cap'n Refsmmat, Carbuncle, Carlw4514, Chris Dybala, Chris Rodgers, Cosmium, Crash Underride, Crystal whacker, Daniel bg, Darrien, David Gerard, Decumanus, DeltaT, Devil Master, Devleenasamanta, Dissident, Djinn112, Doczilla, DopefishJustin, Drjezza, Drslopey, Edgar181, Emperorbma, Eric119, Femto, Flying Jazz, Fonzy, Fvw, Gotthic, Gveret Tered, Hashar, Heracletus, Hoaxe, Hugo-cs, Hydrargyrum, Icairns, Icarus3, Idiotsjoy, Jerrywills, Joanjoc, John, Jonathan Drain, Jose Ramos, Joshmaul, Kalaong, Keenan Pepper, Kelovy, Kingdon, Kurykh, Kwamikagami, Lengau, Luna Santin, Materialscientist, Mav, Mcsee, Merovingian, Michael Hardy, Mihai, Mike Rosoft, Mike4ty4, Mild Bill Hiccup, Moogsi, Mooncow, MrJones, Muéro, NEMT, Nergaal, Netdragon, Noisy, Nozzleman, Oghmoir, Ortolan88, Ostrich11, Pakaran, Pascal666, Poolkris, Pras, Professor Ninja, RODERICKMOLASAR, Rathlan, Reddi, Remember, Rich Farmbrough, RickK, Rickjpelleg, Rifleman 82, Rjwilmsi, Roberta F., Roentgenium111, Romanskolduns, Rpf2019, Rursus, SakotGrimshine, Saperaud, Schneelocke, Scog, SenseOnes, ShaunMacPherson, Shawis, Shimgray, SierraSciSPA, SimonP, Siroxo, Sl, Snowdog, SolarWind, Stone, Strait, Tagishsimon, Tempodivalse, Tonyrex, Trollminator, Ununseptium, Vicki Rosenzweig, VmanBG, Warut, Yekrats, Yerpo, Zerathzul, 148 anonymous edits
Image Sources, Licenses and Contributors image:Uup-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Uup-TableImage.png License: GNU Free Documentation License Contributors: Conscious, Kwamikagami, Saperaud File:Speaker Icon.svg Source: http://en.wikipedia.org/w/index.php?title=File:Speaker_Icon.svg License: Public Domain Contributors: Blast, G.Hagedorn, Mobius, 2 anonymous edits Image:Yes check.svg Source: http://en.wikipedia.org/w/index.php?title=File:Yes_check.svg License: GNU Free Documentation License Contributors: User:Gmaxwell, User:WarX Image:X mark.svg Source: http://en.wikipedia.org/w/index.php?title=File:X_mark.svg License: GNU Free Documentation License Contributors: Abnormaal, Gmaxwell, Kilom691, MGA73, Mardetanha, Penubag, Pseudomoi, WikipediaMaster, 1 anonymous edits Image:Black x.svg Source: http://en.wikipedia.org/w/index.php?title=File:Black_x.svg License: unknown Contributors: User:Howcheng
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6
Ununhexium
1
Ununhexium ununpentium ← ununhexium → ununseptium Po ↑ Uuh ↓ (Uhh) Periodic Table Extended Periodic Table General Name, symbol, number
ununhexium, Uuh, 116
Element category
presumably poor metal
Group, Period, Block
16, 7, p
Standard atomic weight
−1
[293] g·mol
Electron configuration
14
10
2
4
perhaps [Rn] 5f 6d 7s 7p (guess based on polonium)
Electrons per shell
2, 8, 18, 32, 32, 18, 6
CAS registry number
54100-71-9 Most-stable isotopes
Main article: Isotopes of ununhexium iso
NA
half-life
293
syn
61 ms
292
syn
18 ms
291
syn
290
syn
Uuh Uuh Uuh Uuh
DM α
DE (MeV)
DP
10.54
289
α
10.66
288
18 ms
α
10.74
287
7.1 ms
α
10.84
286
Uuq Uuq Uuq Uuq
References
Ununhexium (pronounced /ˌjuːnənˈhɛksiəm/;[1] officially, the two initial u's are to be [2] pronounced English pronunciation: /uː/ ( listen) ) is the temporary name of a synthetic superheavy element in the periodic table that has the temporary symbol Uuh and has the atomic number 116. About 30 atoms of ununhexium have been produced to date, either directly or as a decay product of ununoctium, and are associated with decays from the four neighbouring isotopes with masses 290-293. The most stable isotope to date is Uuh-293 with a half-life of 63 ms.
Ununhexium
De Facto Discovery On July 19, 2000, scientists at Dubna (FLNR) detected a single decay from an atom of ununhexium following the irradiation of a Cm-248 target with Ca-48 ions. The results were published in December, 2000.[3] This 10.54 MeV alpha-emitting activity was originally assigned to 292Uuh due to the correlation of the daughter to previously assigned 288Uuq. However, that assignment was later altered to 289Uuq, and hence this activity was correspondingly changed to 293Uuh. Two further atoms were reported by the institute during their second experiment between April-May 2001.[4]
In the same experiment they also detected a decay chain which corresponded to the first observed decay of ununquadium and assigned to 289Uuq.[4] This activity has not been observed again in a repeat of the same reaction. However, its detection in this series of experiments indicates the possibility of the decay of an isomer of ununhexium, namely 293b 116, or a rare decay branch of the already discovered isomer,293a116, in which the first alpha particle was missed. Further research is required to positively assign this activity. The team repeated the experiment in April-May 2005 and detected 8 atoms of ununhexium. The measured decay data confirmed the assignment of the discovery isotope as 293116. In this run, the team also observed 292116 in the 4n channel for the first time. [5] In May 2009, the JWP reported on the discovery of element 112 (see ununbium) and acknowledged the discovery of the isotope 283112.[6] This implies the de facto discovery of element 116, as 291116 (see below), from the acknowledgment of the data relating to the granddaughter 283112, although the actual discovery experiment may be determined as that above. An impending JWP report will discuss these issues further.
Naming The element with Z=116 is historically known as eka-polonium. Ununhexium (Uuh) is a temporary IUPAC systematic element name. Research scientists usually refer to the element simply as element 116 (E116).
Proposed names by claimants According to IUPAC recommendations, the discoverer(s) of a new element has the right to suggest a name.[7] The JWP has not yet officially accepted the discovery of element 116 and so the naming process has not yet begun.
Extrapolated chemical properties of eka-polonium Oxidation states Element 116 is projected to be the fourth member of the 7p series of non-metals and the heaviest member of group 16 (VIA) in the Periodic Table, below polonium. The group oxidation state of +VI is known for all the members apart from oxygen which lacks available d-orbitals for expansion and is limited to a maximum +II state, exhibited in the fluoride OF2. The +IV is known for sulfur, selenium, tellurium, and polonium, undergoing a shift in stability from reducing for S(IV) and Se(IV) to oxidizing in Po(IV). Tellurium(IV) is the most stable for this element. This suggests a decreasing stability for the higher oxidation states as the group is descended and element 116 should portray an oxidizing +IV
2
Ununhexium state and a more stable +II state. The lighter members are also known to form a −II state as oxide, sulfide, selenide, and telluride. Polonide formation is unconfirmed or only transient. The extrapolated electronegativity of ununhexium should eliminate this low oxidation state.
Chemistry The possible chemistry of element 116 can be extrapolated from that of polonium. It should therefore undergo oxidation to a dioxide, UuhO2, although a trioxide, UuhO3 is plausible, but unlikely. The stability of a +II state should manifest itself in the formation of a simple monoxide, UuhO. Fluorination will likely result in a tetrafluoride, UuhF4 and/or a difluoride, UuhF2. Chlorination and bromination may well stop at the corresponding dihalides, UuhCl2 and UuhBr2. Oxidation by iodine should certainly stop at UuhI2 and may even be inert to this element.
History of synthesis of isotopes by cold fusion 208
Pb(82Se,xn)290−x116
In 1998, the team at GSI attempted the synthesis of 290116 as a radiative capture (x=0) product. No atoms were detected providing a cross section limit of 4.8 pb.
History of synthesis of isotopes by hot fusion This section deals with the synthesis of nuclei of ununhexium by so-called "hot" fusion reactions. These are processes which create compound nuclei at high excitation energy (~40-50 MeV, hence "hot"), leading to a reduced probability of survival from fission. The excited nucleus then decays to the ground state via the emission of 3-5 neutrons. Fusion reactions utilizing 48Ca nuclei usually produce compound nuclei with intermediate excitation energies (~30-35 MeV) and are sometimes referred to as "warm" fusion reactions. This leads, in part, to relatively high yields from these reactions. 238
U(54Cr,xn)292−x116
There are sketchy indications that this reaction was attempted by the team at GSI in 2006. There are no published results on the outcome, presumably indicating that no atoms were detected. This is expected from a study of the systematics of cross sections for U-238 targets.[8] 248
Cm(48Ca,xn)296−x116 (x=3,4)
The first attempt to synthesise element 116 was performed in 1977 by Ken Hulet and his team at the Lawrence Livermore National Laboratory (LLNL). They were unable to detect any atoms of ununhexium.[9] Yuri Oganessian and his team at the Flerov Laboratory of Nuclear Reactions (FLNR) subsequently attempted the reaction in 1978 and were met by failure. In 1985, a joint experiment between Berkeley and Peter Armbruster's team at GSI, the result was again negative with a calculated cross-section limit of 10–100 pb.[10] In 2000, Russian scientists at Dubna finally succeeded in detecting a single atom of element 116, assigned to the isotope 292116.[3] In 2001, they repeated the reaction and formed a further 2 atoms in a confirmation of their discovery experiment. A third atom was
3
Ununhexium
4
tentatively assigned to 293116 on the basis of a missed parental alpha decay.[4] In April 2004, the team ran the experiment again at higher energy and were able to detect a new decay chain, assigned to 292116. On this basis, the original data were reassigned to 293116. The tentative chain is therefore possibly associated with a rare decay branch of this isotope. 293 [5] In this reaction, 2 further atoms of 116 were detected. 245
Cm(48Ca,xn)293−x116 (x=2,3)
In order to assist in the assignment of isotope mass numbers for ununhexium, in March-May 2003 the Dubna team bombarded a Cm-245 target with Ca-48 ions. They were able to observe two new isotopes, assigned to 291116 and 290116.[11] This experiment was successfully repeated in Feb-March 2005 where 10 atoms were created with identical decay data to those reported in the 2003 experiment.[12]
Synthesis of ununhexium as a decay product Ununhexium has also been observed in the decay of ununoctium. In October 2006 it was announced that 3 atoms of ununoctium had been detected by the bombardment of californium-249 with calcium-48 ions, which then rapidly decayed into ununhexium.[12] The observation of 290116 allowed the assignment of the product to synthesis of a nucleus with Z=118 (see ununoctium).
294
118 and proved the
Chronology of isotope discovery Isotope
Year discovered
Discovery reaction [13] Cf(48Ca,3n)
290
2002
249
291
2003
245
[11]
292
2004
248
[5]
293
2000
248
[3]
Uuh Uuh Uuh Uuh
Cm(48Ca,2n) Cm(48Ca,4n) Cm(48Ca,3n)
Yields of isotopes Hot fusion The table below provides cross-sections and excitation energies for hot fusion reactions producing ununhexium isotopes directly. Data in bold represent maxima derived from excitation function measurements. + represents an observed exit channel. Projectile
Target
CN
48
248
296
48
245
293
Ca Ca
Cm Cm
2n
3n [5] 1.1 pb, 38.9 MeV
Uuh Uuh
4n
[11]
0.9 pb, 33.0 MeV
[11]
3.7 pb, 37.9 MeV
[5] 3.3 pb, 38.9 MeV
5n
Ununhexium
5
Retracted isotopes 289
116
In 1999, researchers at Lawrence Berkeley National Laboratory announced the synthesis of 293 118 (see ununoctium), in a paper published in Physical Review Letters.[14] The claimed isotope 289116 decayed by 11.63 MeV alpha emission with a half-life of 0.64 ms. The following year, they published a retraction after other researchers were unable to duplicate the results.[15] In June 2002, the director of the lab announced that the original claim of the discovery of these two elements had been based on data fabricated by the principal author Victor Ninov. As such, this isotope of E116 is currently unknown.
Fission of compound nuclei with Z=116 Several experiments have been performed between 2000-2006 at the Flerov laboratory of Nuclear Reactions in Dubna studying the fission characteristics of the compound nuclei 296,294,290 116. Four nuclear reactions have been used, namely 248Cm+48Ca, 246Ca+48Ca, 244 50 Pu+ Ti and 232Th+58Fe. The results have revealed how nuclei such as this fission predominantly by expelling closed shell nuclei such as 132Sn (Z=50, N=82). It was also found that the yield for the fusion-fission pathway was similar between 48Ca and 58Fe projectiles, indicating a possible future use of 58Fe projectiles in superheavy element formation.In addition, in comparative experiments synthesizing 294116 using 48Ca and 50Ti projectiles, the yield from fusion-fission was ~3x less for 50Ti, also suggesting a future use in SHE production[16]
Theoretical calculations on decay characteristics Theoretical calculation in a quantum tunneling model supports the experimental data relating to the synthesis of 293,292116.[17] [18]
Target-Projectile Combinations leading to Z=116 compound nuclei The below table contains various combinations of targets and projectiles which could be used to form compound nuclei with Z=116. A indicates a successful reaction, a indicates a failure to date and a indicates a reaction yet to be attempted.
Target-projectile combinations leading to Z=116 nuclei Target
Projectile
CN
208
Pb
82
Se
290
232
Th
58
Fe
290
238
116 116
U
54
244
Pu
50
Ti
294
248
Cm
48
Ca
296
246
48
294
Cm
Cr
Ca
292
116 116 116 116
Attempted?
Ununhexium
6
245
48
Cm
293
Ca
249
116
40
Cf
289
Ar
116
Theoretical calculations on evaporation cross sections The below table contains various targets-projectile combinations for which calculations have provided estimates for cross section yields from various neutron evaporation channels. The channel with the highest expected yield is given. DNS = Di-nuclear system ; σ = cross section Target
Projectile
CN
Channel (product)
~σ
max
Model
Ref
208
82
290
1n (289116)
0.1 pb
DNS
[19]
208
79
287
1n (286116)
0.5 pb
DNS
[19]
Pb Pb
Se Se
116 116
U
54
Cr
292
2n (290116)
0.1 pb
DNS
[20]
250
Cm
48
Ca
298
4n (294116)
5 pb
DNS
[20]
248
48
296
4n (292116)
2 pb
DNS
[20]
247
48
295
3n (292116)
3 pb
DNS
[20]
245
48
293
3n (290116)
1.5 pb
DNS
[20]
238
Cm Cm Cm
Ca Ca Ca
116 116 116 116 116
Future experiments The team at Dubna have indicated plans to synthesize element 116 using the reaction between plutonium-244 and titanium-50. This experiment will allow them to assess the feasibility of using projectiles with Z>20 required in the synthesis of superheavy elements with Z>118. Although initially scheduled for 2008, the reaction looking at the synthesis of evaporation residues has not been conducted to date.[21] There are also plans to repeat the Cm-248 reaction at different projectile energies in order to probe the 2n channel, leading to the new isotope 294116. In addition, they have future plans to complete the excitation function of the 4n channel product, 292116, which will allow them to assess the stabilizing effect of the N=184 shell on the yield of evaporation residues. The GSI also have plans to study the formation of 293,292116 in the 248Cm(48Ca,xn) reaction as a first step in their future program with a 248Cm target, aiming towards a synthesis of element 120 (see unbinilium).
Ununhexium
7
See also • Island of stability
External links • WebElements.com: Uuh [22] • Apsidium: Ununhexium 116 [23] • Second postcard from the island of stability
[24]
References [1] (http:/ / reference. aol. com/ columbia/ _a/ ununhexium/ 20051207161909990011) [2] http:/ / media. iupac. org/ publications/ pac/ 1979/ pdf/ 5102x0381. pdf [3] "Observation of the decay of 292116" (http:/ / prola. aps. org/ abstract/ PRC/ v63/ i1/ e011301), Oganessian et al., Phys. Rev. C 63, 011301 (2000). Retrieved 2008-03-03 [4] "Confirmed results of the 248Cm(48Ca,4n)292116 experiment" (https:/ / e-reports-ext. llnl. gov/ pdf/ 302186. pdf), Patin et al., LLNL report (2003). Retrieved 2008-03-03 [5] "Measurements of cross sections and decay properties of the isotopes of elements 112, 114, and 116 produced in the fusion reactions 233,238U , 242Pu , and 248Cm+48Ca" (http:/ / prola. aps. org/ abstract/ PRC/ v70/ i6/ e064609), Oganessian et al., Phys. Rev. C 70, 064609 (2004). Retrieved 2008-03-03 [6] R.C.Barber; H.W.Gaeggeler;P.J.Karol;H. Nakahara; E.Verdaci; E. Vogt (2009). " Discovery of the element with atomic number 112 (http:/ / media. iupac. org/ publications/ pac/ asap/ pdf/ PAC-REP-08-03-05. pdf)" (IUPAC Technical Report). Pure Appl. Chem.. doi: 10.1351/PAC-REP-08-03-05 (http:/ / dx. doi. org/ 10. 1351/ PAC-REP-08-03-05). . [7] http:/ / media. iupac. org/ publications/ pac/ 2002/ pdf/ 7405x0787. pdf [8] "List of experiments 2000-2006" (http:/ / opal. dnp. fmph. uniba. sk/ ~beer/ experiments. php) [9] " Search for Superheavy Elements in the Bombardment of 248Cm with 48Ca (http:/ / prola. aps. org/ abstract/ PRL/ v39/ i7/ p385_1)". . [10] " Attempts to Produce Superheavy Elements by Fusion of 48Ca with 248Cm in the Bombarding Energy Range of 4.5-5.2 MeV/u (http:/ / prola. aps. org/ abstract/ PRL/ v54/ i5/ p406_1)". . [11] "Measurements of cross sections for the fusion-evaporation reactions 244Pu(48Ca,xn)292−x114 and 245 Cm(48Ca,xn)293−x116" (http:/ / prola. aps. org/ abstract/ PRC/ v69/ i5/ e054607), Oganessian et al., Phys. Rev. C 69, 054607 (2004). Retrieved 2008-03-03 [12] "Synthesis of the isotopes of elements 118 and 116 in the 249Cf and 245Cm+48Ca fusion reactions" (http:/ / scitation. aip. org/ getabs/ servlet/ GetabsServlet?prog=normal& id=PRVCAN000074000004044602000001& idtype=cvips& gifs=yes), Oganessian et al., Phys. Rev. C 74, 044602 (2006). Retrieved 2008-03-03 [13] see ununoctium [14] Ninov, V.; et al. (1999). " Observation of Superheavy Nuclei Produced in the Reaction of (http:/ / link. aps. org/ abstract/ PRL/ v83/ p1104)". Physical Review Letters 83: 1104. doi: 10.1103/PhysRevLett.83.1104 (http:/ / dx. doi. org/ 10. 1103/ PhysRevLett. 83. 1104). .
86
Kr with
208
Pb
[15] Editorial note on the preceding (http:/ / link. aps. org/ abstract/ PRL/ v89/ e039901). [16] see Flerov lab annual reports 2000-2006 inclusive http:/ / www1. jinr. ru/ Reports/ Reports_eng_arh. html [17] P. Roy Chowdhury, C. Samanta, and D. N. Basu (26 January 2006). " α decay half-lives of new superheavy elements (http:/ / scitation. aip. org/ getabs/ servlet/ GetabsServlet?prog=normal& id=PRVCAN000073000001014612000001& idtype=cvips& gifs=yes)". Phys. Rev. C 73: 014612. doi: 10.1103/PhysRevC.73.014612 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 73. 014612). . Retrieved 2008-01-18. [18] C. Samanta, P. Roy Chowdhury and D.N. Basu (2007). " Predictions of alpha decay half lives of heavy and superheavy elements (http:/ / www. sciencedirect. com/ science?_ob=ArticleURL& _udi=B6TVB-4NF4F0Y-2& _user=2806701& _rdoc=1& _fmt=& _orig=search& _sort=d& view=c& _acct=C000058844& _version=1& _urlVersion=0& _userid=2806701& md5=3f680654b5659191d67f31681a4cfc83)". Nucl. Phys. A 789: 142–154. doi: 10.1016/j.nuclphysa.2007.04.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2007. 04. 001). . [19] http:/ / arxiv. org/ pdf/ 0707. 2588 [20] http:/ / arxiv. org/ pdf/ 0803. 1117 [21] http:/ / flerovlab. jinr. ru/ flnr/ programme_synth_2008. html [22] http:/ / webelements. com/ ununhexium/ [23] http:/ / apsidium. com/ elements/ 116. htm [24] http:/ / cerncourier. com/ main/ article/ 41/ 8/ 17
Article Sources and Contributors
Article Sources and Contributors Ununhexium Source: http://en.wikipedia.org/w/index.php?oldid=307820553 Contributors: Adamfinmo, Ahoerstemeier, Alexius08, Alfio, AlimanRuna, Bcorr, Benbest, BlueEarth, Bobblewik, Brett Dunbar, Bryan Derksen, Buuuuuh, CERminator, Cacahueten, Carbuncle, Ch'marr, Circeus, ClickRick, Closedmouth, Cmprince, Cosmium, Crystal whacker, Darrien, David Gerard, Dbenbenn, Devleenasamanta, Discospinster, Drjezza, Edcolins, Edgar181, Emperorbma, Eric119, Femto, Flying Jazz, Fonzy, Fvw, Gazno, Gertlex, Hashar, Heracletus, Hillman, Hmmm, Hugo-cs, Icairns, Iwoelbern, Jak123, James086, Joanjoc, John, Johnlogic, Jpgordon, Keenan Pepper, Kelovy, Ken Gallager, Khukri, Kingdon, KittySaturn, Kurykh, Kwamikagami, L33tminion, Magnus.de, Materialscientist, Mav, Mazca, Melaen, Merovingian, Michael93555, MightyWarrior, Mike Rosoft, Mpatel, Müslimix, Nergaal, Noisy, Oghmoir, Oppertunity, Ortolan88, Pakaran, Pascal666, Patrick jones, Pawl Kennedy, Peter, Phil Boswell, PoliteCarbide, Poolkris, Publunch, Rathlan, Remember, Rich Farmbrough, RickK, Rickjpelleg, Rifleman 82, Roentgenium111, Romanskolduns, Rursus, Sakinho, Saperaud, Schmerf, ShaunMacPherson, Shimgray, SimonP, Siroxo, Skizzik, SkyLined, Sl, Squids and Chips, Srleffler, Stevey7788, Stirling Newberry, Stone, TJ Spyke, Tagishsimon, Timc, Tonyrex, Trollminator, Vicki Rosenzweig, Yekrats, Zinc90, 116 anonymous edits
Image Sources, Licenses and Contributors image:Uuh-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Uuh-TableImage.png License: GNU Free Documentation License Contributors: Kwamikagami, Paddy, Saperaud, רדמ לבוי File:Speaker Icon.svg Source: http://en.wikipedia.org/w/index.php?title=File:Speaker_Icon.svg License: Public Domain Contributors: Blast, G.Hagedorn, Mobius, 2 anonymous edits Image:Yes check.svg Source: http://en.wikipedia.org/w/index.php?title=File:Yes_check.svg License: GNU Free Documentation License Contributors: User:Gmaxwell, User:WarX Image:X mark.svg Source: http://en.wikipedia.org/w/index.php?title=File:X_mark.svg License: GNU Free Documentation License Contributors: Abnormaal, Gmaxwell, Kilom691, MGA73, Mardetanha, Penubag, Pseudomoi, WikipediaMaster, 1 anonymous edits Image:Black x.svg Source: http://en.wikipedia.org/w/index.php?title=File:Black_x.svg License: unknown Contributors: User:Howcheng
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8
Ununseptium
1
Ununseptium ununhexium ← Ununseptium → ununoctium At ↑ Uus ↓ (Uhs) Periodic Table Extended Periodic Table General Name, symbol, number Element category Group, Period, Block Appearance
Ununseptium, Uus, 117 Unknown 17, 7, p Unknown
Standard atomic weight Electron configuration Electrons per shell Phase
−1
Unknown g·mol Unknown
2,8,18,32,32,18,7 Unknown
CAS registry number
87658-56-8 References
Ununseptium (pronounced /ˌjuːnənˈsɛptiəm/; officially, the two initial u's are to be [1] pronounced English pronunciation: /uː/ ( listen) ) is the temporary name of an undiscovered chemical element in the periodic table that has the temporary symbol Uus and the atomic number 117. It is the only missing element in period 7 of the periodic table. Since it is placed below the halogens it may share qualities similar to astatine or iodine. The first attempt to synthesise this element is currently underway at the Flerov Laboratory of Nuclear Reactions in Dubna, Russia.
Naming The element with Z=117 is historically known as eka-astatine (see 'eka' terminology). The name ununseptium is a systematic element name, used as a placeholder until the element is discovered, the discovery is acknowledged by the IUPAC, and the IUPAC decides on a name. Usually, the name suggested by the discoverer(s) is chosen. According to current guidelines from IUPAC, the ultimate name for element 117 must end in -ium, not -ine, even if ununseptium turns out to be a halogen.[2]
Ununseptium
2
Current experiments The team at the Flerov laboratory of nuclear reactions has begun an experiment to synthesize element 117 using the reaction below:[3] [4]
The expected cross-section is of the order of ~2 pb. The expected evaporation residues, 293 117 and 294117, are predicted to decay via relatively long decay chains as far as isotopes of dubnium or lawrencium.
Calculated decay chains from the parent nuclei and 294 Uus
293
Uus
url=http:/ / 159. 93. 28. 88/ linkc/ education/ SHE_Sagaidak. pdf|accessdate=2009-07-07}}
Target-Projectile Combinations leading to Z=117 compound nuclei The below table contains various combinations of targets and projectiles which could be used to form compound nuclei with Z=117. A indicates a successful reaction, a indicates a failure to date and a indicates a reaction yet to be attempted.
Target-projectile combinations leading to Z=117 nuclei Target
Projectile
CN
208
Pb
81
Br
289
232
Th
59
Co
291
U
55
Mn
293
Np
54
238
237
Cr
117 117 117
291
117
244
Pu
51
V
295
243
Am
50
Ti
293
248
Cm
45
Sc
293
48
Ca
297
249
Bk
117 117 117 117
Attempted?
Ununseptium
3
249
41
Cf
290
K
117
Theoretical calculations on evaporation cross sections The below table contains various targets-projectile combinations for which calculations have provided estimates for cross section yields from various neutron evaporation channels. The channel with the highest expected yield is given. DNS = Di-nuclear system ; σ = cross section Target
Projectile
CN
Channel (product)
~σ
max
Model
Ref
209
82
291
1n (290117)
15 fb
DNS
[5]
209
Bi
79
Se
288
1n (287117)
0.2 pb
DNS
[5]
232
Th
59
Co
291
2n (289117)
0.1 pb
DNS
[5]
55
Mn
293
2-3n (291,290117)
70 fb
DNS
[5]
V
295
3n (292117)
0.6 pb
DNS
[5]
Sc
Bi
238
U
244
Pu
Se
51
117 117 117 117 117
248
45
293
4n (289117)
2.9 pb
DNS
[5]
246
45
Sc
291
4n (287117)
1 pb
DNS
[5]
249
48
Ca
297
3n (294117)
2.1 pb ; 3 pb
DNS
[5] [6]
247
48
295
3n (292117)
0.8 pb ; 0.9 pb
DNS
[6] [5]
Cm Cm Bk Bk
Ca
117 117 117 117
Future experiments The team at the GSI in Darmstadt, recently acknowledged as the discoverers of element 112 (see ununbium) have begun experiments aiming towards a synthesis of element 117. The GSI have indicated that if they are unable to acquire any 249Bk from the United States, which is likely given the situation regarding the attempt in Russia, they will study the reaction 244Pu(51V,xn) instead, or possibly 243Am(50Ti,xn).[7]
Predicted Decay Characteristics Theoretical calculations in a quantum tunneling model with mass estimates from a macroscopic-microscopic model predict the alpha-decay half-lives of isotopes of the element 117 (namely, 289-303117) to be around 0.1-40 ms.[8] [9] [10]
Predicted chemical properties Certain chemical properties, such as bond lengths, are predicted to differ from what one would expect based on periodic trends from the lighter halogens (because of relativistic effects).[11]
Ununseptium
References [1] [2] [3] [4] [5] [6] [7] [8]
http:/ / media. iupac. org/ publications/ pac/ 1979/ pdf/ 5102x0381. pdf http:/ / media. iupac. org/ publications/ pac/ 2002/ pdf/ 7405x0787. pdf http:/ / rss. russiatoday. ru/ Top_News/ 2009-07-01/ New_chemical_element_to_be_synthesized_in_Russia. html http:/ / flerovlab. jinr. ru/ flnr/ experiments. html http:/ / arxiv. org/ pdf/ 0708. 0159 http:/ / arxiv. org/ pdf/ 0803. 1117 http:/ / www-win. gsi. de/ tasca08/ contributions/ TASCA08_Cont_Duellmann2b. pdf C. Samanta, P. Roy Chowdhury and D.N. Basu (2007). "Predictions of alpha decay half lives of heavy and superheavy elements". Nucl. Phys. A 789: 142. doi: 10.1016/j.nuclphysa.2007.04.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2007. 04. 001). [9] P. Roy Chowdhury, C. Samanta, and D. N. Basu (2008). "Search for long lived heaviest nuclei beyond the valley of stability". Phys. Rev. C 77: 044603. doi: 10.1103/PhysRevC.77.044603 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 77. 044603). [10] P. Roy Chowdhury, C. Samanta, and D. N. Basu (2008). "Nuclear half-lives for α -radioactivity of elements with 100 ≤ Z ≤ 130". At. Data & Nucl. Data Tables 94: 781-806. doi: 10.1016/j.adt.2008.01.003 (http:/ / dx. doi. org/ 10. 1016/ j. adt. 2008. 01. 003). [11] Trond Saue. " Principles and Applications of Relativistic Molecular Calculations (http:/ / dirac. chem. sdu. dk/ thesis/ 96. saue_phd. pdf)". ., page 76
4
Article Sources and Contributors
Article Sources and Contributors Ununseptium Source: http://en.wikipedia.org/w/index.php?oldid=307909370 Contributors: 7'ad, Ahoerstemeier, AlimanRuna, Allen3, Armando12, Astavats, Bhangranuch, BlueEarth, Bryan Derksen, CBDunkerson, Cacahueten, Calvero JP, Camembert, Carbuncle, Cax17, ChongDae, Cosmium, Cromwellt, Daniel Quinlan, Darrien, David Latapie, Dbeaglefan2, Discospinster, Dlohcierekim's sock, Drjezza, Edgar181, EmilJ, Emperorbma, Encyclopedia77, Epbr123, Eric119, Femto, FoekeNoppert, Fonzy, Fvw, Giftlite, GregorB, Grika, Gveret Tered, H264, Halcatalyst, Hashar, Hugo-cs, Isnow, Jehan60188, Jess Cully, Joanjoc, John, KarlFrei, Kay Dekker, Kelovy, Kingdon, KnightRider, Kwamikagami, Louis Waweru, Materialscientist, Mav, Merovingian, Mike Rosoft, Mollsmolyneux, Nergaal, Newone, Oghmoir, Ortolan88, Pascal666, PeterJeremy, Pgk, Pi399, Pill, PoliteCarbide, Poolkris, Rathlan, Remember, RickK, Rickjpelleg, Roentgenium111, Romanskolduns, Samir, Saperaud, Sengkang, SimonP, Siroxo, Skizzik, Sl, SlamDiego, Stagerj, Super-Magician, Superheavy120, The way, the truth, and the light, TheMadBaron, Tide rolls, Trollminator, UtherSRG, Vicki Rosenzweig, Yekrats, 102 anonymous edits
Image Sources, Licenses and Contributors image:Uus-TableImage.png Source: http://en.wikipedia.org/w/index.php?title=File:Uus-TableImage.png License: GNU Free Documentation License Contributors: Kwamikagami, Saperaud File:Speaker Icon.svg Source: http://en.wikipedia.org/w/index.php?title=File:Speaker_Icon.svg License: Public Domain Contributors: Blast, G.Hagedorn, Mobius, 2 anonymous edits Image:293Uus and 294Uus calculated decay chains.jpg Source: http://en.wikipedia.org/w/index.php?title=File:293Uus_and_294Uus_calculated_decay_chains.jpg License: Creative Commons Zero Contributors: drjezza Image:249Bk+48Ca calculated excitation function.jpg Source: http://en.wikipedia.org/w/index.php?title=File:249Bk+48Ca_calculated_excitation_function.jpg License: Free Art License Contributors: drjezza Image:Yes check.svg Source: http://en.wikipedia.org/w/index.php?title=File:Yes_check.svg License: GNU Free Documentation License Contributors: User:Gmaxwell, User:WarX Image:X mark.svg Source: http://en.wikipedia.org/w/index.php?title=File:X_mark.svg License: GNU Free Documentation License Contributors: Abnormaal, Gmaxwell, Kilom691, MGA73, Mardetanha, Penubag, Pseudomoi, WikipediaMaster, 1 anonymous edits Image:Black x.svg Source: http://en.wikipedia.org/w/index.php?title=File:Black_x.svg License: unknown Contributors: User:Howcheng
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5
Ununoctium
1
Ununoctium ununseptium ← ununoctium → ununenniumRn ↑ Uuo ↓ (Uho)
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HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygenFluorineNeonSodiumMagnesiumAlumini (element)ThalliumLeadBismuthPoloniumAstatineRadonFranciumRadiumActiniumThoriumProtactiniumU
118Uuo Periodic table
Appearance Unknown General Name, symbol, numberElement categoryGroup, period, blockStandard atomic weightElectron configurationElectrons per shell Physical properties PhaseDensity (near r.t.) Boiling pointCritical pointHeat of fusionHeat of vaporization Atomic properties Oxidation states Ionization energies 2nd: (extrapolated) 1450[1] kJ·mol−1Atomic radiusCovalent radius Miscellaneous Crystal structureCAS registry number Most stable isotopes Main article: Isotopes of ununoctium iso 294
Uuo
[2]
N.A. syn
half-life ~0.89 ms
DM α
DE (MeV) 11.65 ± 0.06
DP 290
Uuh
ununoctium, Uuo, 118 Unknown18, 7, p(294) g·mol−1 [Rn] 5f14 6d10 7s2 7p6[3] 2, 8, 18, 32, 32, 18, 8[3] (Image) Unknown (predicted) 13.65[4] g·cm−3 (extrapolated) 320–380[3] K,50–110 °C,120–220 °F (extrapolated) 439[5] K, 6.8[5] MPa (extrapolated) 23.5[5] kJ·mol−1 (extrapolated) 19.4[5] kJ·mol−1 0[6] , +2[7] , +4[7] 1st: (calc.) 820–1130[3] kJ·mol−1 (predicted) 152[4] pm (extrapolated) 230[1] pm Unknown 54144-19-3[8] Ununoctium (pronounced /ˌjuːnəˈnɒktiəm/;[9] officially, the two initial Us are to be [10] pronounced English pronunciation: /uː/ ( listen) ), also known as eka-radon or element 118, is the temporary IUPAC name[11] for the transactinide element having the atomic number 118 and temporary element symbol Uuo. On the periodic table of the elements, it is a p-block element and the last one of the 7th period. Ununoctium is currently the only
Ununoctium
2
synthetic member of group 18. It has the highest atomic number and highest atomic mass of all discovered elements. The radioactive ununoctium atom is very unstable, and since 2002, only three atoms (possibly four) of the isotope 294Uuo have been detected.[12] While this allowed for very little experimental characterization of its properties and possible compounds, theoretical calculations have allowed for many predictions, including some very unexpected ones. For example, although ununoctium is a member of group 18, it is probably not a noble gas like all the other group 18 elements.[3] It was formerly thought to be a gas but is now predicted to be a solid under normal conditions.[3]
History Unsuccessful attempts In late 1998, Polish physicist Robert Smolanczuk published calculations on the fusion of atomic nuclei towards the synthesis of superheavy atoms, including element 118.[13] His calculations suggested that it might be possible to make element 118 by fusing lead with krypton under carefully controlled conditions.[13] In 1999, researchers at Lawrence Berkeley National Laboratory made use of these predictions and announced the discovery of elements 116 and 118, in a paper published in Physical Review Letters,[14] and very soon after the results were reported in Science.[15] The researchers claimed to have performed the reaction: 8636Kr
+
20882Pb
→
293118Uuo
+n
The following year, they published a retraction after researchers at other laboratories were unable to duplicate the results and the Berkeley lab itself was unable to duplicate them as well.[16] In June 2002, the director of the lab announced that the original claim of the discovery of these two elements had been based on data fabricated by principal author Victor Ninov.[17]
Discovery First decay of atoms of ununoctium was observed at the JINR in Dubna in 2002.[18] On October 9, 2006, researchers from Joint Institute for Nuclear Research (JINR) and Lawrence Livermore National Laboratory of California, USA, working at the JINR in Dubna, Russia, announced in Physical Review C[2] that they had indirectly detected a total of three (possibly four) nuclei of ununoctium-294 (one or two in 2002[19] and two more in 2005) produced via collisions of californium-249 atoms and calcium-48 ions:[20] [21] [22] [23] [24] 24998Cf
+
4820Ca
→
294118Uuo
+3n
Ununoctium
3 Because of the very small fusion reaction probability (the fusion cross section is ~0.3-0.6 pb = (3-6)×10−41 m2) the experiment took 4 months and involved a beam dose of 4×1019 calcium ions that had to be shot at the californium target to produce the first recorded event believed to be the synthesis of ununoctium.[8] Nevertheless, researchers are highly confident that the results are not a false positive, since the chance that the detections were random events was estimated to be less than one part in 100,000.[25]
Radioactive decay pathway of isotope [2] Uuo-294. The decay energy and average half-life is given for the parent isotope and each daughter isotopes. The fraction of atoms undergoing spontaneous fission (SF) is given in green.
In the experiments, the alpha-decay of three atoms of ununoctium was observed. A fourth decay by direct spontaneous fission was also proposed. A half-life of 0.89 ms was calculated: 294Uuo decays into 290Uuh by alpha decay. Since there were only three nuclei, the halflife derived from observed lifetimes has a large uncertainty:
0.89(+1.07|-0.31)ms.[2] 294118Uuo
→
290116Uuh
+ 4He
The identification of the 294Uuo nuclei was verified by separately creating the putative daughter nucleus 290Uuh by means of a bombardment of 245Cm with 48Ca ions, 24596Cm
+
4820Ca
→
290116Uuh
+3n
and checking that the 290Uuh decay matched the decay chain of the 294Uuo nuclei.[2] The daughter nucleus 290Uuh is very unstable, decaying with a halflife of 14 milliseconds into 286 Uuq, which may undergo spontaneous fission or alpha decay into 282Uub, which will undergo spontaneous fission.[2] In a quantum-tunneling model, the alpha decay half-life of 294118 was predicted to be 0.66(+0.23,-0.18)ms[26] with the experimental Q-value published in 2004.[27] Calculation with theoretical Q-values from the macroscopic-microscopic model of Muntian-Hofman-Patyk-Sobiczewski gives somewhat low but comparable results.[28] Following the success in obtaining ununoctium, the discoverers have started similar experiments in the hope of creating element 120 from 58Fe and 244Pu.[29] Isotopes of the element 120 are predicted to have alpha decay half lives of the order of micro-seconds.[30] [31]
Ununoctium
Naming Until the 1960s ununoctium was known as eka-emanation (emanation is the old name for radon).[32] In 1979 the IUPAC published recommendations according to which the element was to be called ununoctium,[33] a systematic element name, as a placeholder until the discovery of the element is confirmed and the IUPAC decides on a name. Before the retraction in 2002, the researchers from Berkeley had intended to name the element ghiorsium (Gh), after Albert Ghiorso (a leading member of the research team).[34] The Russian discoverers reported their synthesis in 2006. In 2007, the head of the Russian institute stated the team were considering two names for the new element: Flyorium in honor of Georgy Flyorov, the founder of the research institute; and moskovium, in recognition of the Moskovskaya Oblast where Dubna is located.[35] He also stated that although the element was discovered as an American collaboration, who provided the californium target, the element should rightly be named in honour of Russia since the Flerov Laboratory of Nuclear Reactions at JINR was the only facility in the world which could achieve this result.[36] [37]
Predicted Characteristics Nucleus stability and isotopes There are no elements with an atomic number above 83 (after bismuth) that have stable isotopes (or measurable decay lifetimes). The stability of nuclei decreases with the increase in atomic number, such that all isotopes with an atomic number above 101 decay radioactively with a half-life under a day. Nevertheless, due to Element 118 comes right at the end of the "island of reasons not very well understood yet, there stability" and thus its nuclei are slightly more stable than otherwise predicted. is a slight increased nuclear stability around elements 110–114, which leads to the appearance of what is known in nuclear physics as the "island of stability". This concept, proposed by UC Berkeley professor Glenn Seaborg, explains why superheavy elements last longer than predicted.[38] Ununoctium is radioactive and has half-life that appears to be less than a millisecond. Nonetheless, this is still longer than some predicted values,[26] [39] thus giving further support to the idea of this "island of stability".[40] Calculations using a quantum-tunneling model predict the existence of several neutron-rich isotopes of element 118 with alpha-decay half-lives close to 1 ms.[30] [31] Theoretical calculations done on the synthetic pathways for, and the half-life of, other isotopes have shown that some could be slightly more stable than the synthesized isotope 294 Uuo, most likely 293Uuo, 295Uuo, 296Uuo, 297Uuo, 298Uuo, 300Uuo and 302Uuo.[26] [41] Of these, 297Uuo might provide the best chances for obtaining longer-lived nuclei,[26] [41] and thus might become the focus of future work with this element. Some isotopes with many more neutrons, such as some located around 313Uuo, could also provide longer-lived nuclei.[42]
4
Ununoctium
5
Calculated atomic and physical properties Ununoctium is a member of group 18, the zero-valence elements. The members of this group are usually inert to most common chemical reactions (for example, combustion) because the outer valence shell is completely filled with eight electrons. This produces a stable, minimum energy configuration in which the outer electrons are tightly bound.[43] It is thought that similarly, ununoctium has a closed outer valence shell in which its valence electrons are arranged in a 7s2, 7p6 configuration.[3] Consequently, some expect ununoctium to have similar physical and chemical properties to other members of its group, most closely resembling the noble gas above it in the periodic table, radon.[44] Following the periodic trend, ununoctium would be expected to be slightly more reactive than radon. However, theoretical calculations have shown that it could be quite reactive, so that it can probably not be considered a noble gas.[7] In addition to being far more reactive than radon, ununoctium may be even more reactive than elements 114 and 112.[3] The reason for the apparent enhancement of the chemical activity of element 118 relative to radon is an energetic destabilization and a radial expansion of the last occupied 7p subshell.[3] [45] More precisely, considerable spin-orbit interactions between the 7p electrons with the inert 7s2 electrons, effectively lead to a second valence shell closing at element 114, and a significant decrease in stabilization of the closed shell of element 118.[3] It has also been calculated that ununoctium, unlike other noble gases, binds an electron with release of energy—or in other words, it exhibits positive electron affinity.[46] [47] [48] Ununoctium is expected to have by far the broadest polarizability of all elements before it in the periodic table, and almost twofold of radon.[3] By extrapolating from the other noble gases, it is expected that ununoctium has a boiling point between 320 and 380 K.[3] This is very different from the previously estimated values of 263 K[1] or 247 K.[49] Even given the large uncertainties of the calculations, it seems highly unlikely that element 118 would be a gas under standard conditions.[3] [50] And as the liquid range of the other gases is very narrow, between 2 and 9 kelvins, this element should be solid at standard conditions. If ununoctium forms a gas under standard conditions nevertheless, it would be one of the densest substances gaseous at standard conditions (even if it is monatomic like the other noble gases). Because of its tremendous polarizability, ununoctium is expected to have an anomalously low ionization energy (similar to that of lead which is 70% of that of radon[51] and significantly smaller than that of element 114[52] ) and a standard state condensed phase.[3]
Predicated compounds No compounds of ununoctium have been synthesized yet, but calculations on theoretical compounds have been performed since 1964.[32] It is expected that if the ionization energy of the element is high enough, it will be difficult to oxidize and therefore, the most common oxidation state will be 0 (as for other noble gases).[6] Calculations on the dimeric molecule Uuo2 showed a bonding interaction roughly equivalent to that calculated for Hg2, and a
XeF4 and RnF4 have a square planar configuration
Ununoctium
6
dissociation energy of 6 kJ/mol, roughly 4 times of that of Rn2.[3] But most strikingly, it was calculated to have a bond length shorter than in Rn2 by 0.16 Å, which would be indicative of a significant bonding interaction.[3] On the other hand, the compound UuoH+ exhibits a dissociation energy (in other words proton affinity of Uuo) that is smaller than that of RnH+.[3] The bonding between ununoctium and hydrogen in UuoH is very weak and can be regarded as a pure van der Waals interaction rather than a true chemical bond.[51] On the other hand, with UuoF4 is predicted to have highly electronegative elements, ununoctium seems to form a tetrahedral configuration more stable compounds than for example element 112 or element 114.[51] The stable oxidation states +2 and +4 have been predicted to exist in the fluorinated compounds UuoF2 and UuoF4.[53] This is a result of the same spin-orbit interactions that make ununoctium unusually reactive. For example, it was shown that the reaction of Uuo with F2 to form the compound UuoF2, would release an energy of 106 kcal/mol of which about 46 kcal/mol come from these interactions.[51] For comparison, the spin-orbit interaction for the similar molecule RnF2 is about 10 kcal/mol out of a formation energy of 49 kcal/mol.[51] The same interaction stabilizes the tetrahedral Td configuration for UuoF4, as opposed to the square planar D4h one of XeF4 and RnF4.[53] The Uuo-F bond will most probably be ionic rather than covalent, rendering the UuoFn compounds non-volatile.[7] [54] Unlike the other noble gases, ununoctium was predicted to be sufficiently electropositive to form a Uuo-Cl bond with chlorine.[7]
Target-Projectile Combinations leading to Z=118 compound nuclei The below table contains various combinations of targets and projectiles which could be used to form compound nuclei with Z=118. A indicates a successful reaction, a indicates a failure to date and a indicates a reaction yet to be attempted.
Target-projectile combinations leading to Z=118 nuclei Target
Projectile
CN
208
Pb
86
232
Th
64
Ni
296
U
58
Fe
296
244
Pu
54
248
Cm
50
Ti
298
48
Ca
297
238
249
Cf
Kr
Cr
294
118 118 118
298
118 118 118
Attempted?
Ununoctium
7
Theoretical calculations on evaporation cross sections The below table contains various targets-projectile combinations for which calculations have provided estimates for cross section yields from various neutron evaporation channels. The channel with the highest expected yield is given. DNS = Di-nuclear system ; σ = cross section Target
Projectile
CN
Channel (product)
~σ
max
Model
Ref
208
86
294
1n (293118)
0.1 pb
DNS
[55]
208
85
Kr
293
1n (292118)
0.18 pb
DNS
[55]
252
48
Ca
300
3n (297118)
1.2 pb
DNS
[56]
251
48
299
3n (296118)
1.2 pb
DNS
[56]
249
48
297
3n (294118)
0.3 pb
DNS
[56]
Pb Pb Cf Cf Cf
Kr
Ca Ca
118 118 118 118 118
Future uses Since only four atoms of ununoctium have ever been produced, it currently has no uses outside of basic scientific research. It would constitute a radiation hazard if enough were ever assembled in one place.[57]
See also • Transuranic element • Ununhexium
External links • ELEMENT 118: EXPERIMENTS on DISCOVERY [58], archive of discoverers' official web page • Chemistry-Blog: Independent analysis of 118 claim [59] • WebElements: Ununoctium [60] • Apsidium: Ununoctium - Moskowium [61] • It's Elemental: Ununoctium [62] • On the Claims for Discovery of Elements 110, 111, 112, 114, 116, and 118 (IUPAC Technical Report) [63]
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8
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J. Phys. G: Nucl. Part. Phys. 34: R165–R242. doi: 10.1088/0954-3899/34/4/R01 (http:/ / dx. doi. org/ 10. 1088/ 0954-3899/ 34/ 4/ R01). [40] " New Element Isolated Only Briefly (http:/ / www. dailycal. org/ printable. php?id=21871)". The Daily Californian. 2006-10-18. . Retrieved 2008-01-18. [41] G. Royer, K. Zbiri, C. Bonilla (2004). "Entrance channels and alpha decay half-lives of the heaviest elements". Nuclear Physics A 730: 355–376. doi: 10.1016/j.nuclphysa.2003.11.010 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2003. 11. 010). [42] S B Duarte, O A P Tavares, M Gonçalves, O Rodríguez, F Guzmán, T N Barbosa, F García and A Dimarco (2004). "Half-life predictions for decay modes of superheavy nuclei". J. Phys. G: Nucl. Part. Phys. 30: 1487–1494. doi: 10.1088/0954-3899/30/10/014 (http:/ / dx. doi. org/ 10. 1088/ 0954-3899/ 30/ 10/ 014). [43] Bader, Richard F.W.. " An Introduction to the Electronic Structure of Atoms and Molecules (http:/ / miranda. chemistry. mcmaster. ca/ esam/ )". McMaster University. . Retrieved 2008-01-18. [44] " Ununoctium (Uuo) - Chemical properties, Health and Environmental effects (http:/ / lenntech. com/ Periodic-chart-elements/ Uuo-en. htm)". Lenntech. . Retrieved 2008-01-18. [45] the actual quote is: "The reason for the apparent enhancement of chemical activity of element 118 relative to radon is the energetic destabilization and radial expansion of its occupied 7p3/2 spinor shell" [46] Igor Goidenko, Leonti Labzowsky, Ephraim Eliav, Uzi Kaldor, and Pekka Pyykko¨ (2003). "QED corrections to the binding energy of the eka-radon (Z=118) negative ion". Physical Review A 67: 020102(R). doi: 10.1103/PhysRevA.67.020102 (http:/ / dx. doi. org/ 10. 1103/ PhysRevA. 67. 020102). [47] Ephraim Eliav and Uzi Kaldor (1996). "Element 118: The First Rare Gas with an Electron Affinity". Physical Review Letters 77 (27): 5350. doi: 10.1103/PhysRevLett.77.5350 (http:/ / dx. doi. org/ 10. 1103/ PhysRevLett. 77. 5350). [48] Nevertheless, quantum electrodynamic corrections have been shown to be quite significant in reducing this affinity (by decreasing the binding in the anion Uuo− by 9%) thus confirming the importance of these corrections in superheavy atoms. See Pyykko [49] N. Takahashi (2002). "Boiling points of the superheavy elements 117 and 118". Journal of Radioanalytical and Nuclear Chemistry 251 (2): 299–301. doi: 10.1023/A:1014880730282 (http:/ / dx. doi. org/ 10. 1023/ A:1014880730282).
9
Ununoctium [50] It is debatable if the name of the group 'noble gases' will be changed if ununoctium is shown to be non-volatile. [51] Young-Kyu Han, Cheolbeom Bae, Sang-Kil Son, and Yoon Sup Lee (2000). "Spin–orbit effects on the transactinide p-block element monohydrides MH (M=element 113–118)". Journal of Chemical Physics 112 (6). doi: 10.1063/1.480842 (http:/ / dx. doi. org/ 10. 1063/ 1. 480842). [52] Clinton S. Nash (1999). "Spin-Orbit Effects, VSEPR Theory, and the Electronic Structures of Heavy and Superheavy Group IVA Hydrides and Group VIIIA Tetrafluorides. A Partial Role Reversal for Elements 114 and 118". J. Phys. Chem. A 1999 (3): 402–410. doi: 10.1021/jp982735k (http:/ / dx. doi. org/ 10. 1021/ jp982735k). [53] Young-Kyu Han and Yoon Sup Lee (1999). "Structures of RgFn (Rg = Xe, Rn, and Element 118. n = 2, 4.) Calculated by Two-component Spin-Orbit Methods. A Spin-Orbit Induced Isomer of (118)F4". J. Phys. Chem. A 103 (8): 1104–1108. doi: 10.1021/jp983665k (http:/ / dx. doi. org/ 10. 1021/ jp983665k). [54] Kenneth S. Pitzer (1975). "Fluorides of radon and element 118". J. Chem. Soc., Chem. Commun.: 760b–761. doi: 10.1039/C3975000760b (http:/ / dx. doi. org/ 10. 1039/ C3975000760b). [55] http:/ / arxiv. org/ pdf/ 0707. 2588 [56] http:/ / arxiv. org/ pdf/ 0803. 1117 [57] " Ununoctium: Biological information (http:/ / webelements. com/ webelements/ elements/ text/ Uuo/ biol. html)". WebElements Periodic Table. . Retrieved 2008-01-18. [58] http:/ / web. archive. org/ web/ 20061129112314/ http:/ / flerovlab. jinr. ru/ flnr/ elm118. html [59] http:/ / www. chemistry-blog. com/ 2006/ 10/ 16/ discovery-of-element-118-by-oganessian-dont-call-it-ununoctium/ [60] http:/ / webelements. com/ ununoctium/ [61] http:/ / www. apsidium. com/ elements/ 118. htm [62] http:/ / education. jlab. org/ itselemental/ ele118. html [63] http:/ / iupac. org/ publications/ pac/ 75/ 10/ 1601/
10
Article Sources and Contributors
Article Sources and Contributors Ununoctium Source: http://en.wikipedia.org/w/index.php?oldid=309087209 Contributors: Act333, Adashiel, Addshore, Ahoerstemeier, Aja-Oki, Alerante, AlimanRuna, Almit39, Altenmann, Anarchist42, Andres, Angstorm, Anthony Appleyard, Atlant, Atrian, Baccyak4H, Ben Tillman, Benbest, Bhangranuch, Black and White, Blindman shady, BlueEarth, Bmdavll, Borgx, Boris Barowski, Brian Kendig, Brighterorange, Brouhaha, Bryan Derksen, COGDEN, Cacahueten, Camembert, Can't sleep, clown will eat me, CaptainMike, Ch'marr, Chris Dybala, Chymicus, CieloEstrellado, Cmapm, Cosmium, Crazy coyote, D Monack, DDima, Daniel bg, Danny, Darrien, David Haslam, David Latapie, DerHexer, Devleenasamanta, Dionyziz, Dismas, Dispenser, Drjezza, Duphus2, Dysprosia, EdJohnston, Eddideigel, Edgar181, El C, Eleassar, Emperorbma, Eric119, Error411, Expensivehat, Fabiform, Femto, Fonzy, Frankenpuppy, Frotz, Furrykef, Fvw, GTBacchus, GW Simulations, GaiusTimiusAwesomus, Gamma, Gerweck, Golthar, Gorman, Gotyear, Grandmasterka, GregorB, Grimhelm, Gveret Tered, Hans van Deukeren, Happy-melon, Hashar, Hateless, HenkvD, Henrygb, Herbee, Heron, Honeycake, Hugo-cs, Hules001, Iamthewalrus36, Iluvcapra, Isnow, Itub, Iwoelbern, Ixfd64, JLM, Jared Hunt, Jess Cully, Jiang, Jimaginator, Joanjoc, John, Julius Sahara, KFan II, Kingdon, Kurykh, Kwamikagami, Laur, Leapmark, Lesgles, Lewis R, Lexicon, Lightmouse, Lumos3, Lysdexia, MER-C, Mac Davis, Mallanox, Malleus Fatuorum, Man123123, Materialscientist, Matthew kokai, Mav, Mcsee, Merovingian, Micahbrwn, Michael Devore, Michael Hardy, Michał Sobkowski, Mike Rosoft, Millar153, Minghong, Miranche, Miss Madeline, Mitchandre, Mouse is back, Mpatel, Mr. Lefty, Mushroom, Mxn, NERIC-Security, Natox, Negrulio, Neitherday, NerdyNSK, Nergaal, Nightstallion, Nilfanion, Noisy, Nozzleman, NuclearVacuum, NuclearWarfare, Numbo3, Obey, Oghmoir, Olin, Omicronpersei8, One, Ortolan88, Oscar Bravo, Oscarcwk, PJY, Pampas Cat, Pascal666, Physicsdavid, Pifreak94, PoliteCarbide, Poolkris, Prozacchiwawa, QuantumEngineer, RJHall, RJaguar3, Randomblue, Rathlan, Raven4x4x, Rawr, Rbraunwa, Repaxan, Reywas92, Rich Farmbrough, RickK, Rickjpelleg, Rjwilmsi, Rmhermen, RobertG, Roentgenium111, RoryReloaded, RoyBoy, Ryomaandres, Saperaud, Schneelocke, Sengkang, Seth Ilys, Shalom Yechiel, Shandris, Shimgray, Short Verses, SimonP, Siroxo, Sk4p, SkyLined, Sl, Socram, Spellage, SteveRwanda, Stevebeck, Strait, Sturmde, Styrofoam1994, Super Rad!, Synthetic element, THEN WHO WAS PHONE?, Tarret, TheMadBaron, TheNewPhobia, Thingg, Timneu22, Timwi, Treisijs, Trollminator, Twinxor, Vicki Rosenzweig, Vuo, Wachholder0, Wafulz, Wimvandorst, Woosie roxs, Xaosflux, Xoloz, Yamaguchi先生, Yekrats, Zfr, Zimbardo Cookie Experiment, Zzyzx11, 234 anonymous edits
Image Sources, Licenses and Contributors file:Unknown.svg Source: http://en.wikipedia.org/w/index.php?title=File:Unknown.svg License: Public Domain Contributors: Mav file:Electron shell 118 Ununoctium.svg Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_118_Ununoctium.svg License: Creative Commons Attribution-Sharealike 2.0 Contributors: User:GregRobson, User:Pumbaa80 File:Speaker Icon.svg Source: http://en.wikipedia.org/w/index.php?title=File:Speaker_Icon.svg License: Public Domain Contributors: Blast, G.Hagedorn, Mobius, 2 anonymous edits Image:Ununoctium-294 nuclear.png Source: http://en.wikipedia.org/w/index.php?title=File:Ununoctium-294_nuclear.png License: unknown Contributors: Original uploader was Nergaal at en.wikipedia Image:Island-of-Stability.png Source: http://en.wikipedia.org/w/index.php?title=File:Island-of-Stability.png License: GNU Free Documentation License Contributors: Original uploader was Xanthine at en.wikipedia Later version(s) were uploaded by McLoaf at en.wikipedia. Image:Square-planar-3D-balls.png Source: http://en.wikipedia.org/w/index.php?title=File:Square-planar-3D-balls.png License: Public Domain Contributors: Benjah-bmm27, Zzyzx11 Image:Tetrahedral-3D-balls.png Source: http://en.wikipedia.org/w/index.php?title=File:Tetrahedral-3D-balls.png License: Public Domain Contributors: Benjah-bmm27 Image:Yes check.svg Source: http://en.wikipedia.org/w/index.php?title=File:Yes_check.svg License: GNU Free Documentation License Contributors: User:Gmaxwell, User:WarX Image:X mark.svg Source: http://en.wikipedia.org/w/index.php?title=File:X_mark.svg License: GNU Free Documentation License Contributors: Abnormaal, Gmaxwell, Kilom691, MGA73, Mardetanha, Penubag, Pseudomoi, WikipediaMaster, 1 anonymous edits Image:Black x.svg Source: http://en.wikipedia.org/w/index.php?title=File:Black_x.svg License: unknown Contributors: User:Howcheng
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
11
Ununennium
1
Ununennium 119
ununoctium ← Ununennium → unbinilium
Fr ↑
Uue ↓
Uhe periodic table - extended periodic table
General Name, Symbol, Number
Ununennium, Uue, 119
Element category
Unknown
Group, Period, Block
1, 8, s
Appearance
Unknown Standard atomic weight
−1
Unknown g·mol
Electron configuration
1
[Uuo] 8s
Electrons per shell
2, 8, 18, 32, 32, 18, 8, 1 Physical properties
Phase
Unknown
Miscellaneous Most-stable isotopes Main article: Isotopes of ununennium iso
NA
half-life
DM
DE (MeV)
DP
References
Ununennium (pronounced /ˌjuːnəˈnɛniəm/; officially, the two initial u's are to be [1] pronounced English pronunciation: /uː/ ( listen) ), or eka-francium, is the temporary name of a hypothetical chemical element in the periodic table that has the temporary symbol Uue and has the atomic number 119. Since it is below the alkali metals it might have properties similar to those of francium or cesium. Like other alkali metals, it should be extremely reactive with water and air. Ununennium would be the first element in the eighth period of the periodic table and the seventh alkali metal.
Ununennium
2
Unsuccessful attempts at synthesis The synthesis of element 119 was attempted in 1985 by bombarding a target of einsteinium-254 with calcium-48 ions at the superHILAC accelerator at Berkeley, California. No atoms were identified, leading to a limiting yield of 300 nb.[2]
It is highly unlikely that this reaction will be useful given the extremely difficult task of making sufficient amounts of Es-254 to make a large enough target to increase the sensitivity of the experiment to the required level, due to the rarity of the element, and extreme rarity of the isotope.
Predicted decay characteristics The alpha-decay half-lives of 1700 nuclei with 100 ≤ Z ≤ 130 have been calculated in a quantum tunneling model with alpha-decay Q-values from different mass estimates.[3] [4] [5] The alpha-decay half-lives predicted for 291-307119 are of the order of micro-seconds. The highest value of the alpha-decay half-life predicted in the quantum tunneling model with the mass estimates from a macroscopic-microscopic model is ~485 microseconds for the isotope 294119. For 302119 it is ~163 microseconds.
Target-Projectile Combinations leading to Z=119 compound nuclei The below table contains various combinations of targets and projectiles which could be used to form compound nuclei with Z=119. A indicates a successful reaction, a indicates a failure to date and a indicates a reaction yet to be attempted.
Target-projectile combinations leading to Z=119 nuclei Target
Projectile
CN
208
Pb
87
Rb
295
232
Th
65
Cu
297
238
119 119
U
59
Np
58
Fe
295
244
Pu
55
Mn
299
243
Am
54
237
248
Co
Cr
297
119 119 119
297
119
51
V
299
50
Ti
299
249
Cf
45
Sc
294
254
Es
48
Ca
302
Cm
249
Bk
119 119 119 119
Attempted?
Ununennium
3
Theoretical calculations on evaporation cross sections The below table contains various targets-projectile combinations for which calculations have provided estimates for cross section yields from various neutron evaporation channels. The channel with the highest expected yield is given. DNS = Di-nuclear system ; σ = cross section Target 254
Es
Projectile 48
Ca
CN 302
119
Channel (product) 3n (299119)
~σ
max
0.5 pb
Model
Ref
DNS
[6]
See also • ununoctium — unbinilium
References [1] http:/ / media. iupac. org/ publications/ pac/ 1979/ pdf/ 5102x0381. pdf [2] R. W. Lougheed, J. H. Landrum, E. K. Hulet, J. F. Wild, R. J. Dougan, A. D. Dougan, H. Gäggeler, M. Schädel, K. J. Moody, K. E. Gregorich, and G. T. Seaborg (1985). " Search for superheavy elements using 48Ca + 254Esg reaction (http:/ / link. aps. org/ abstract/ PRC/ v32/ p1760)". Physical Reviews C 32: 1760–1763. doi: 10.1103/PhysRevC.32.1760 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 32. 1760). . [3] C. Samanta, P. Roy Chowdhury and D.N. Basu (2007). "Predictions of alpha decay half lives of heavy and superheavy elements". Nucl. Phys. A 789: 142–154. doi: 10.1016/j.nuclphysa.2007.04.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2007. 04. 001). [4] P. Roy Chowdhury, C. Samanta, and D. N. Basu (2008). "Search for long lived heaviest nuclei beyond the valley of stability". Phys. Rev. C 77: 044603. doi: 10.1103/PhysRevC.77.044603 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 77. 044603). [5] P. Roy Chowdhury, C. Samanta, and D. N. Basu (2008). "Nuclear half-lives for α -radioactivity of elements with 100 ≤ Z ≤ 130". At. Data & Nucl. Data Tables 94: 781–806. doi: 10.1016/j.adt.2008.01.003 (http:/ / dx. doi. org/ 10. 1016/ j. adt. 2008. 01. 003). [6] http:/ / arxiv. org/ pdf/ 0803. 1117
Article Sources and Contributors
Article Sources and Contributors Ununennium Source: http://en.wikipedia.org/w/index.php?oldid=307909763 Contributors: Allen3, Amalas, Anthony Appleyard, Aranherunar, Bhangranuch, BlueEarth, Bueller 007, Cax17, Cfailde, Chillysnow, Complex01, Cosmium, Darrien, David Latapie, Db099221, Dcorrin, DoubleBlue, Drjezza, E23, Edgar181, Emperorbma, Eric119, Femto, Feneeth of Borg, Fiveless, Fonzy, Frencheigh, Gertlex, Headbomb, Honeycake, Hopfer25148, INVERTED, Ixfd64, Jab416171, Jagro, John, JohnyDog, Karlchwe, Kate, Keenan Pepper, Kelovy, Kurykh, Kwamikagami, Larry laptop, Lysdexia, Materialscientist, Mav, Mike Rosoft, Miranche, NawlinWiki, Nergaal, NickBush24, Nimbulan, Oppertunity, Ortolan88, Poolkris, Roentgenium111, RoryReloaded, Scienceguy2005, Srinivasasha, Stone, Su-no-G, Superheavy120, TheRabidMonkie, Timc, Tomwalden, Trixt, Warut, Wikianon, Yewlongbow, 66 anonymous edits
Image Sources, Licenses and Contributors Image:-TableImage.svg Source: http://en.wikipedia.org/w/index.php?title=File:-TableImage.svg License: GNU Free Documentation License Contributors: Bastique, Bibi Saint-Pol, Kwamikagami, Lantrix, Mdd4696, Popolon, Soeb, Tietew, 1 anonymous edits File:Speaker Icon.svg Source: http://en.wikipedia.org/w/index.php?title=File:Speaker_Icon.svg License: Public Domain Contributors: Blast, G.Hagedorn, Mobius, 2 anonymous edits Image:Yes check.svg Source: http://en.wikipedia.org/w/index.php?title=File:Yes_check.svg License: GNU Free Documentation License Contributors: User:Gmaxwell, User:WarX Image:X mark.svg Source: http://en.wikipedia.org/w/index.php?title=File:X_mark.svg License: GNU Free Documentation License Contributors: Abnormaal, Gmaxwell, Kilom691, MGA73, Mardetanha, Penubag, Pseudomoi, WikipediaMaster, 1 anonymous edits Image:Black x.svg Source: http://en.wikipedia.org/w/index.php?title=File:Black_x.svg License: unknown Contributors: User:Howcheng
License Creative Commons Attribution-Share Alike 3.0 Unported http:/ / creativecommons. org/ licenses/ by-sa/ 3. 0/
4
Unbinilium
1
Unbinilium 120
Uue ← unbinilium → Ubu
Ra ↑
Ubn ↓
Usn periodic table - extended periodic table
General Name, Symbol, Number
unbinilium, Ubn, 120
Element category
Presumably Alkali earth metals
Group, Period, Block
2, 8, s
Appearance
dark metallic ? Standard atomic weight
−1
unknown g·mol
Electron configuration
2
[Uuo] 8s (a guess based upon barium and radium)
Electrons per shell
2, 8, 18, 32, 32, 18, 8, 2 Physical properties
Oxidation states
presumably 2
Phase
presumably a solid
Miscellaneous Most-stable isotopes Main article: Isotopes of unbinilium iso
NA
half-life
DM
DE (MeV)
DP
References [1] Unbinilium (pronounced /ˌuːnbaɪˈnɪliəm/ ( listen) ), also called eka-radium, is the temporary, systematic element name of a hypothetical chemical element in the periodic table that has the temporary symbol Ubn and has the atomic number 120. Since unbinilium is placed below the alkaline earth metals it possibly has properties similar to those of radium or barium.
Attempts to date to synthesize the element using fusion reactions at low excitation energy has met with failure although there are reports that the fission of nuclei of element 120 at very high excitation has been successfully measured, indicating a strong shell effect at Z=120.
Unbinilium
Attempts at synthesis Neutron evaporation In March-April 2007, the synthesis of element 120 was attempted at the Flerov Laboratory of Nuclear Reactions in Dubna by bombarding a plutonium-244 target with iron-58 ions.[2] Initial analysis revealed that no atoms of element 120 were produced providing a limit of 400 fb for the cross section at the energy studied.[3]
The Russian team is planning to upgrade their facilities before attempting the reaction again. In April 2007, the team at GSI attempted to create unbinilium using uranium-238 and nickel-64:[4]
No atoms were detected providing a limit of 1.6 pb on the cross section at the energy provided. The GSI repeated the experiment with higher sensitivity in three separate runs from April-May 2007, Jan- March 2008 and Sept- Oct 2008, all with negative results and providing a cross section limit of 90 fb.[5]
Compound nucleus fission Element 120 is of interest because it is part of the hypothesized island of stability, with the compound nucleus 302120 being the most stable of those that can be created directly by current methods. It has been calculated that Z=120 may in fact be the next proton magic number, rather than at Z=114 or 126. Several experiments have been performed between 2000-2008 at the Flerov laboratory of Nuclear Reactions in Dubna studying the fission characteristics of the compound nucleus 302 120. Two nuclear reactions have been used, namely 244Pu+58Fe and 238U+64Ni. The results have revealed how nuclei such as this fission predominantly by expelling closed shell nuclei such as 132Sn (Z=50, N=82). It was also found that the yield for the fusion-fission pathway was similar between 48Ca and 58Fe projectiles, indicating a possible future use of 58Fe projectiles in superheavy element formation.[6] In 2008, the team at GANIL, France, described the results from a new technique which attempts to measure the fission half-life of a compound nucleus at high excitation energy, since the yields are significantly higher than from neutron evaporation channels. It is also a useful method for probing the effects of shell closures on the survivability of compound nuclei in the super-heavy region, which can indicate the exact position of the next proton shell (Z=114, 120, 124 or 126). The team studied the nuclear fusion reaction between uranium ions and a target of natural nickel:
The results indicated that nuclei of element 120 were produced at high (~70 MeV) excitation energy which underwent fission with measurable half-lives > 10-18s. Although very short, the ability to measure such a process indicates a strong shell effect at Z=120. At lower excitation energy (see neutron evaporation), the effect of the shell will be enhanced and ground-state nuclei can be expected to have relatively long half-lives. This result could partially explain the relatively long half-life of 294118 measured in experiments at Dubna (see ununoctium). Similar experiments have indicated a similar phenomenon at Z=124 (see
2
Unbinilium
3
unbiquadium) but not at Z = 114 (see ununquadium), suggesting that the next proton shell does in fact lie at Z>120.[7] [8]
Future Reactions The GSI have plans to start up a program utilizing 248Cm targets for SHE production and will most likely attempt this reaction in 2010 or 2011.[9] Likewise, the team at RIKEN have also begun a program utilizing 248Cm targets and have indicated future experiments to probe the possibility of Z=120 being the next magic number using the aforementioned nuclear reactions to form 302120.[10]
Calculated Decay Characteristics In a quantum tunneling model with mass estimates from a macroscopic-microscopic model, the alpha-decay half-lives of several isotopes of the element 120 (namely, 292-304120) have been predicted to be around 1-20 microseconds. [11] [12] [13] [14]
Extrapolated Reactivity Unbinilium should be highly reactive, according to periodic trends, as this element is a member of alkaline earth metals. It would be much more reactive than any other lighter elements of this group. If group reactivity is followed, this element would react violently in air to form an oxide (UbnO) and in water to form the hydroxide, which would be a strong base and highly explosive in terms of flammability. It is also possible that, due to relativistic effects, the element has noble gas character, as already seen for element 114.
Target-Projectile Combinations leading to Z=120 Compound nuclei The below table contains various combinations of targets and projectiles which could be used to form compound nuclei with Z=120. A indicates a successful reaction, a indicates a failure to date and a indicates a reaction yet to be attempted.
Target-projectile combinations leading to Z=120 nuclei Target 232
Th
238
Projectile 70
Zn
CN 302
120
U
64
Ni
302
244
Pu
58
Fe
302
248
Cm
54
249
Cf
Cr
50
Ti
120 120
302
120
299
120
Attempted?
Unbinilium
4
Theoretical calculations on evaporation cross sections The below table contains various targets-projectile combinations for which calculations have provided estimates for cross section yields from various neutron evaporation channels. The channel with the highest expected yield is given. DNS = dinuclear system ; AS = advanced statistical ; σ = cross section Target
Projectile
CN
Channel (product)
~σ
max
Model
Ref
208
88
296
1n (295120)
70 fb
DNS
[15]
208
87
Sr
295
1n (294120)
80 fb
DNS
[15]
238
64
Ni
302
2n (300120)
0.5 fb
DNS
[16]
238
U
64
Ni
302
4n (298120)
2 ab
DNS - AS
[17]
244
Pu
58
Fe
302
3n (299120)
8 fb
DNS - AS
[17]
248
Cm
54
302
3n (299120)
10 pb
DNS - AS
[17]
248
54
Cr
302
4n (298120)
30 fb
DNS
[18]
48
Ca
305
3n (302120)
70 fb
DNS
[16]
Pb Pb U
Cm
257
Fm
Sr
Cr
120 120 120 120 120 120 120 120
See also • • • •
Island of stability: Ununquadium – Unbinilium – Unbihexium Radium Eka-francium — Eka-radium — Eka-actinium Ununennium – Unbiunium
External links • WebElements.com - Unbinilium [19] • Apsidium - Unbinilium [20] • Theory of atomic mass calculation [21]
References [1] http:/ / media. iupac. org/ publications/ pac/ 1979/ pdf/ 5102x0381. pdf [2] THEME03-5-1004-94/2009 (http:/ / www. jinr. ru/ plan/ ptp-2007/ e751004. htm) [3] Oganessian et al. (2009). "Attempt to produce element 120 in the 244Pu+58Fe reaction". Phys. Rev. C 79: 024603. doi: 10.1103/PhysRevC.73.014612 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 73. 014612). [4] [5] [6] [7] [8]
http:/ / www. gsi. de/ documents/ DOC-2007-Mar-174-1. pdf http:/ / www. gsi. de/ informationen/ wti/ library/ scientificreport2008/ PAPERS/ NUSTAR-SHE-01. pdf see Flerov lab annual reports 2000-2004 inclusive http:/ / www1. jinr. ru/ Reports/ Reports_eng_arh. html http:/ / physics. aps. org/ articles/ v1/ 12 Phys. Rev. Lett. 101: 072701. http:/ / scitation. aip. org/ getabs/ servlet/ GetabsServlet?prog=normal& id=PRLTAO000101000007072701000001& idtype=cvips& gifs=yes.
[9] http:/ / www-aix. gsi. de/ EA/ (see U248) [10] see slide 11 http:/ / www-win. gsi. de/ tasca07/ contributions/ TASCA07_Contribution_Morita. pdf [11] P. Roy Chowdhury, C. Samanta, and D. N. Basu (26 January 2006). "α decay half-lives of new superheavy elements". Phys. Rev. C 73: 014612. doi: 10.1103/PhysRevC.73.014612 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 73. 014612).
Unbinilium [12] C. Samanta, P. Roy Chowdhury and D.N. Basu (2007). "Predictions of alpha decay half lives of heavy and superheavy elements". Nucl. Phys. A 789: 142–154. doi: 10.1016/j.nuclphysa.2007.04.001 (http:/ / dx. doi. org/ 10. 1016/ j. nuclphysa. 2007. 04. 001). [13] P. Roy Chowdhury, C. Samanta, and D. N. Basu (2008). "Search for long lived heaviest nuclei beyond the valley of stability". Phys. Rev. C 77: 044603. doi: 10.1103/PhysRevC.77.044603 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 77. 044603). [14] P. Roy Chowdhury, C. Samanta, and D. N. Basu (2008). "Nuclear half-lives for α -radioactivity of elements with 100 ≤ Z ≤ 130". At. Data & Nucl. Data Tables 94: 781–806. doi: 10.1016/j.adt.2008.01.003 (http:/ / dx. doi. org/ 10. 1016/ j. adt. 2008. 01. 003). [15] http:/ / arxiv. org/ pdf/ 0707. 2588 [16] http:/ / arxiv. org/ pdf/ 0803. 1117 [17] http:/ / arxiv. org/ pdf/ 0812. 4410 [18] Zagebraev. V ; Greiner, W (2008). "Synthesis of superheavy nuclei: A search for new production reactions". Phys. Rev. C 78: 034610. doi: 10.1103/PhysRevC.78.034610 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 78. 034610). [19] http:/ / www. webelements. com/ webelements/ elements/ text/ Ubn/ index. html [20] http:/ / www. apsidium. com/ elements/ 120. htm [21] http:/ / www. apsidium. com/ number/ amc. htm
5
Article Sources and Contributors
Article Sources and Contributors Unbinilium Source: http://en.wikipedia.org/w/index.php?oldid=307911696 Contributors: Allen3, AussieLegend, Bhangranuch, BlueEarth, Captain panda, Cax17, Charles Matthews, Complex01, Cosmium, Darrien, David Latapie, Drjezza, Eclecticology, Edgar181, Emperorbma, Encyclopedia77, Eric119, Femto, Fonzy, Georgia guy, Gertlex, INVERTED, Jeronimo, John, Kate, Kelovy, Kurykh, Kwamikagami, Materialscientist, Mav, Merovingian, Mike Rosoft, Mikya, Mycroft.Holmes, Nergaal, NetRolller 3D, Nimbulan, Ojs, OldakQuill, Olin, Oneoverzero, Ortolan88, Oscarcwk, Pakaran, PiMaster3, Poolkris, Roo72, RoryReloaded, Skadetdom, Snagglepuss, Son, Stirling Newberry, Superheavy120, Timc, Tooto, Trixt, Van helsing, Vuo, WCFrancis, 41 anonymous edits
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6
Unbiunium
1
Unbiunium 121
Ubn ← unbiunium → Ubb
↑
Ubu ↓
Usu periodic table - extended periodic table
General Name, Symbol, Number
unbiunium, Ubu, 121
Element category
Superactinides
Group, Period, Block
g1, 8, g
Appearance
unknown Standard atomic weight
−1
[320] u (supposition) g·mol
Electron configuration
1
2
[Uuo] 5g 8s
Electrons per shell
2, 8, 18, 32, 33, 18, 8, 2 Physical properties
Phase
presumably solid
Miscellaneous Most-stable isotopes Main article: Isotopes of unbiunium iso
NA
half-life
DM
DE (MeV)
DP
References
Unbiunium (pronounced /ˌuːnbaɪˈjuːniəm/) is the temporary name of a hypothetical chemical element in the periodic table that has the temporary symbol Ubu and has the atomic number 121. Unbiunium is the first element whose ground state electron configuration contains an electron in a g subshell, making it the first element in the g-block. However, neither lanthanum nor actinium show such a predicted ground state for the f-block, and lie in the transition metals, so unbiunium's third electron should also hang up. Other elements have access to their g subshells, though not in their ground states. As of July 2009, no attempt has ever been made to synthesis element 121.
Unbiunium
2
Naming The name unbiunium is a systematic element name, used as a placeholder until it is confirmed by other research groups and the IUPAC decides on a name. Usually, the name suggested by the discoverer(s) is chosen.
Target-Projectile Combinations leading to Z=121 compound nuclei The below table contains various combinations of targets and projectiles which could be used to form compound nuclei with Z=121. A indicates a successful reaction, a indicates a failure to date and a indicates a reaction yet to be attempted.
Target-projectile combinations leading to Z=121 nuclei Target 208
Projectile Y
297
71
Ga
303
U
65
Cu
303
Np
64
Ni
301
244
Pu
59
Co
303
243
Am
58
Fe
301
248
Cm
55
Mn
303
Pb
232
Th
238
237
249
Bk
89
CN
54
Cr
Attempted?
121 121 121 121 121 121 121
303
121
249
Cf
51
V
300
254
Es
50
Ti
304
121 121
External links • Likely properties of Unbiunium [1] • Journal of Chemical Physics, 1998, V 109, N 10, 8 Sep.
See also • Unbinilium–Unbibium
References [1] http:/ / www. apsidium. com/ elements/ 121. htm [2] http:/ / library. iem. ac. ru/ j-ch-ph/ 1998/ 10-10998. html
[2]
Article Sources and Contributors
Article Sources and Contributors Unbiunium Source: http://en.wikipedia.org/w/index.php?oldid=307619587 Contributors: Allen3, Biblbroks, BlueEarth, Cosmium, David Latapie, Dhidalgo, Drjezza, Eric119, Habbit, INVERTED, Jonlys, Keenan Pepper, Kwamikagami, Merovingian, Mintleaf, Nick Y., Nimbulan, Phil Boswell, Poolkris, Roentgenium111, Scepia, Syd Henderson, 30 anonymous edits
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3
Unbibium
1
Unbibium 122
Ubu ← unbibium → Ubt
↑
Ubb ↓
Usb periodic table - extended periodic table
General Name, Symbol, Number
unbibium, Ubb, 122
Element category
Superactinides
Group, Period, Block
g2, 8, g
Appearance
unknown Standard atomic weight
−1
g·mol
Electron configuration
2
2
[Uuo] 5g 8s
Electrons per shell
2, 8, 18, 32, 34, 18, 8, 2 Physical properties
Phase
presumably solid
Miscellaneous Most-stable isotopes Main article: Isotopes of unbibium iso
NA
half-life
DM
DE (MeV)
DP
References
Unbibium (pronounced /uːnˈbaɪbiəm/), also referred to as eka-thorium or simply element 122, is the temporary name of a currently unknown chemical element in the periodic table that has the temporary symbol Ubb and the atomic number 122. In 2008, it was claimed to have been discovered in natural thorium samples[1] but that claim has now been dismissed by recent repetitions of the experiment using more accurate techniques.
Unbibium
2
History Neutron Evaporation The first attempt to synthesize element 122 was performed in 1972 by Flerov et al. at JINR, using the hot fusion reaction:
No atoms were detected and a yield limit of 5 mb (5,000,000 pb) was measured. Current results (see ununquadium) have shown that the sensitivity of this experiment was too low by at least 6 orders of magnitude. In 2000, the Gesellschaft für Schwerionenforschung performed a very similar experiment with much higher sensitivity:
These results indicate that the synthesis of such heavier elements remains a significant challenge and further improvements of beam intensity and experimental efficiency is required. The sensitivity should be increased to 1 fb.
Compound Nucleus Fission Several experiments have been performed between 2000-2004 at the Flerov laboratory of Nuclear Reactions studying the fission characteristics of the compound nucleus 306122. Two nuclear reactions have been used, namely 248Cm+58Fe and 242Pu+64Ni. The results have revealed how nuclei such as this fission predominantly by expelling closed shell nuclei such as 132Sn (Z=50, N=82). It was also found that the yield for the fusion-fission pathway was similar between 48Ca and 58Fe projectiles, indicating a possible future use of 58Fe projectiles in superheavy element formation.[2]
Target-Projectile Combinations leading to Z=122 compound nuclei The below table contains various combinations of targets and projectiles which could be used to form compound nuclei with Z=122. A indicates a successful reaction, a indicates a failure to date and a indicates a reaction yet to be attempted.
Target-projectile combinations leading to Z=122 nuclei Target
Projectile
CN
208
Pb
94
Zr
302
232
Th
74
Ge
306
122 122
238
70
308
238
U
66
304
244
Pu
64
Ni
308
248
Cm
58
Fe
306
U
249
Cf
Zn Zn
54
Cr
122 122 122 122
303
122
Attempted?
Unbibium
3
Claimed discovery as a naturally-occurring element On April 24, 2008, a group led by Amnon Marinov at the Hebrew University of Jerusalem claimed to have found single atoms of unbibium in naturally occurring thorium deposits at an abundance of between 10-11 and 10-12, relative to thorium.[1] The claim of Marinov et al. was criticized by a part of the scientific community, and Marinov says he has submitted the article to the journals Nature and Nature Physics but both turned it down without sending it for peer review.[3] . A criticism of the technique, previously used in purportedly identifying lighter thorium isotopes by mass spectrometry,[4] [5] was published in Physical Review C in 2008.[6] A rebuttal by the Marinov group was published in Physical Review C after the published comment.[7] A repeat of the thorium-experiment using the superior method of Accelerator Mass Spectrometry (AMS) failed to confirm the results, despite a 100x greater sensitivity.[8] This result throws considerable doubt on the results of the Marinov collaboration with regards to their claims of long-lived isotopes of thorium, roentgenium and unbibium.
See also • Island of stability • Systematic element name • Unbiunium – Unbitrium
External links • • • •
Chemistry-Blog: Independent analysis of Marinov’s 122 claim WebElements.com - Unbibium [10] Likely properties of Unbibium [11] Marinov's Site [12]
[9]
References [1] Marinov, A.; Rodushkin, I.; Kolb, D.; Pape, A.; Kashiv, Y.; Brandt, R.; Gentry, R. V.; Miller, H. W. (2008). " Evidence for a long-lived superheavy nucleus with atomic mass number A=292 and atomic number Z=~122 in natural Th (http:/ / arxiv. org/ abs/ 0804. 3869)". arXiv.org. . Retrieved 2008-04-28. [2] see Flerov lab annual reports 2000-2004 inclusive http:/ / www1. jinr. ru/ Reports/ Reports_eng_arh. html [3] Royal Society of Chemistry, Chemistry World, "Heaviest element claim criticised" (http:/ / www. rsc. org/ chemistryworld/ News/ 2008/ May/ 02050802. asp) [4] A. Marinov; I. Rodushkin; Y. Kashiv; L. Halicz; I. Segal; A. Pape; R. V. Gentry; H. W. Miller; D. Kolb; R. Brandt (2007). " Existence of long-lived isomeric states in naturally-occurring neutron-deficient Th isotopes (http:/ / scitation. aip. org/ getabs/ servlet/ GetabsServlet?prog=normal& id=PRVCAN000076000002021303000001& idtype=cvips& gifs=yes)". Phys. Rev. C 76: 021303(R). . [5] http:/ / arxiv. org/ ftp/ nucl-ex/ papers/ 0605/ 0605008. pdf [6] R. C. Barber; J. R. De Laeter (2009). " Comment on “Existence of long-lived isomeric states in naturally-occurring neutron-deficient Th isotopes” (http:/ / scitation. aip. org/ getabs/ servlet/ GetabsServlet?prog=normal& id=PRVCAN000079000004049801000001& idtype=cvips& gifs=yes)". Phys. Rev. C 79: 049801. . [7] A. Marinov; I. Rodushkin; Y. Kashiv; L. Halicz; I. Segal; A. Pape; R. V. Gentry; H. W. Miller; D. Kolb; R. Brandt (2009). " Reply to “Comment on `Existence of long-lived isomeric states in naturally-occurring neutron-deficient Th isotopes'” (http:/ / scitation. aip. org/ getabs/ servlet/ GetabsServlet?prog=normal& id=PRVCAN000079000004049802000001& idtype=cvips& gifs=yes)". Phys. Rev. C 79: 049802. . [8] J. Lachner; I. Dillmann; T. Faestermann; G. Korschinek; M. Poutivtsev; G. Rugel (2008). " Search for long-lived isomeric states in neutron-deficient thorium isotopes (http:/ / scitation. aip. org/ getabs/ servlet/
Unbibium GetabsServlet?prog=normal& id=PRVCAN000078000006064313000001& idtype=cvips& gifs=yes)". Phys. Rev. C 78: 064313. . [9] http:/ / www. chemistry-blog. com/ 2008/ 04/ 29/ adressing-marinovs-element-122-claim/ [10] http:/ / www. webelements. com/ webelements/ elements/ text/ Ubb/ index. html [11] http:/ / www. apsidium. com/ elements/ 122. htm [12] http:/ / www. phys. huji. ac. il/ ~marinov/ index. htm
4
Article Sources and Contributors
Article Sources and Contributors Unbibium Source: http://en.wikipedia.org/w/index.php?oldid=307620342 Contributors: -Midorihana-, Allen3, Biblbroks, BlueEarth, Bogdangiusca, Cosmium, Crakkpot, DMacks, David Latapie, Dor Cohen, Drjezza, Eric119, Femto, Frankchn, Headbomb, Jeltz, John, Jonlys, Kelovy, Kkmurray, Kwamikagami, Michał Sobkowski, Nick Y., Nickj, Nightstallion, Nimbulan, Phil Boswell, Polonium, Poolkris, RJHall, Rursus, Ryant5000, SimonP, Siroxo, SteveBaker, Thue, Tony Sidaway, Warut, Wereon, 40 anonymous edits
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5
Unbitrium
1
Unbitrium 123
Ubb ← unbitrium → Ubq
↑
Ubt ↓
Ust periodic table - extended periodic table
General Name, Symbol, Number
unbitrium, Ubt, 123
Element category
Superactinides
Group, Period, Block
g3, 8, g
Appearance
? Standard atomic weight
[326] g·mol
Electron configuration
[Uuo] 5g 8s
−1
3
Electrons per shell
2
2, 8, 18, 32, 35, 18, 8, 2 Physical properties
Phase
presumably solid
Miscellaneous Most-stable isotopes Main article: Isotopes of unbitrium iso
NA
half-life
DM
DE (MeV)
DP
References
Unbitrium (pronounced /uːnˈbaɪtriəm/) is the temporary name of a hypothetical chemical element in the periodic table that has the temporary symbol Ubt and has the atomic number 123. As of July 2009, no attempt has ever been made to synthesise element 123.
Unbitrium
Naming The name unbitrium is a systematic element name, used as a placeholder until it is confirmed by other research groups and the IUPAC decides on a name. Usually, the name suggested by the discoverer(s) is chosen.
Fictional Reference to Element 123 In the fictional universe of Star Trek: The Next Generation, this element was apparently discovered and named. The episode "Rascals" depicted a "trans-periodic table," in a schoolroom set, which depicted element number 123 as being named jamesium, symbol Rj, having an atomic weight of 326. The graphic was created by set artists, probably a homage to designer Richard D. James, and listed this element in the fictional "gamma series."
External links • WebElements.com - Unbitrium [1] • Likely propeties of Unbitrium [2]
See also • Unbibium – Unbiquadium
References [1] http:/ / www. webelements. com/ webelements/ elements/ text/ Ubt/ index. html [2] http:/ / www. apsidium. com/ elements/ 123. htm
2
Article Sources and Contributors
Article Sources and Contributors Unbitrium Source: http://en.wikipedia.org/w/index.php?oldid=302718846 Contributors: Allen3, AxG, Biblbroks, BlueEarth, CaptainMike, Cosmium, David Latapie, Drjezza, Epbr123, Ephraim33, Eric119, Femto, Headbomb, INVERTED, Kelovy, Kurykh, Kwamikagami, LeafGreen Ranger, Makeemlighter, Nimbulan, Poolkris, Roentgenium111, Shanes, Tavilis, Wtmitchell, 29 anonymous edits
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3
Unbiquadium
1
Unbiquadium 124
Ubt ← unbiquadium → Ubp
↑
Ubq ↓
Usq periodic table - extended periodic table
General Name, Symbol, Number
unbiquadium, Ubq, 124
Element category
Superactinides
Group, Period, Block
g4, 8, g
Appearance
unknown Standard atomic weight
−1
Unknown g·mol
Electron configuration
4
2
[Uuo] 5g 8s
Electrons per shell
2, 8, 18, 32, 36, 18, 8, 2 Physical properties
Phase
Unknown Miscellaneous Most-stable isotopes Main article: Isotopes of unbiquadium
iso
NA
half-life
DM
DE (MeV)
DP
References
Unbiquadium (pronounced /ˌuːnbaɪˈkwɒdiəm/) is the temporary name of a hypothetical element in the periodic table that has the temporary symbol Ubq and atomic number 124. In 2008, a team at GANIL, France, published results indicating that nuclei of element 124 had been produced at very high excitation energy, which underwent fission with measurable lifetimes. This important results suggests a strong stabilizing effect at Z=124 and points to the next proton shell at Z>120, not at Z=114 as previously thought.
Synthesis of Z=124 Nuclei In a series of experiments, scientists at GANIL have attempted to measure delayed fission of compound nuclei of elements with Z=114, 120 and 124 in shell effects in this region and to pinpoint the next spherical proton shell. In results published in 2008, the team provided results from a reaction bombardment of a natural germanium target with uranium ions:
the direct and order to probe 2006, with full involving the
Unbiquadium The team reported that they had been able to identify compound nuclei fissioning with half-lives > 10-18s. Although very short, the ability to measure such decays indicated a strong shell effect at Z=124. A similar phenomenon was found for Z=120 but not for Z=114.[1]
Name The name unbiquadium is an IUPAC systematic element name, the temporary name and symbol assigned to newly-synthesized and not-yet-synthesized chemical elements. A transuranic element receives a permanent name and symbol only after its synthesis has been confirmed. Transuranic elements (those beyond uranium) are, except for microscopic quantities and except for plutonium, always artificially produced, and usually end up being named for a scientist or the location of a laboratory that does work in atomic physics.
See also • Unbitrium−Unbipentium [1] http:/ / hal. archives-ouvertes. fr/ docs/ 00/ 12/ 91/ 31/ PDF/ WAPHE06_EPJ_preprint1. pdf
External links • Webelements.com – Unbiquadium (http:/ / www. webelements. com/ webelements/ elements/ text/ Ubq/ key. html) • Likely properties of unbiquadium (http:/ / www. apsidium. com/ elements/ 124. htm)
2
Article Sources and Contributors
Article Sources and Contributors Unbiquadium Source: http://en.wikipedia.org/w/index.php?oldid=305089852 Contributors: Biblbroks, BlueEarth, Bradeos Graphon, Coredesat, Cosmium, Drbreznjev, Drjezza, Femto, Isaac123456789, Kkmurray, Kurykh, Kwamikagami, Leyo, Roentgenium111, Stirling Newberry, Tardis, Tony Sidaway, Weggie, 10 anonymous edits
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3
Unbipentium
1
Unbipentium 125
Ubq ← unbipentium → Ubh
↑
Ubp ↓
Ust periodic table - extended periodic table
General Name, Symbol, Number
unbipentium, Ubp, 125
Element category
Superactinides
Group, Period, Block
g5, 8, g
Appearance
unknown [[Image: |125px|]]
Standard atomic weight
[332] g·mol
Electron configuration
[Uuo] 5g 8s
−1
5
Electrons per shell
2
2, 8, 18, 32, 37, 18, 8, 2 Physical properties
Phase
presumably solid
Miscellaneous Most-stable isotopes Main article: Isotopes of unbipentium iso
NA
half-life
DM
DE (MeV)
DP
References
Unbipentium (pronounced /ˌuːnbaɪˈpɛntiəm/), or eka-neptunium, is the temporary name of a hypothetical chemical element in the periodic table that has the temporary symbol Ubp and has the atomic number 125. As of July 2009, no attempt has ever been made to synthesise element 125.
Unbipentium
2
Name The name unbipentium is used as a placeholder, such as in scientific articles about the search for element 125. Transuranic elements beyond plutonium are always artificially produced, and usually end up being named for a scientist or the location of a laboratory that does work in atomic physics. Element 125 is of interest because it is within a range that has been predicted to be within a range of experimentally feasible "island of stability" elements based on a mean field theory based on a model of alpha decay.[1] The self-coupling of the ω meson could be responsible for greater shell stability based earlier work, which suggested that self-coupling meson effects were stronger than originally predicted.
See also • Neptunium • Unbiquadium – Unbihexium
External links • Likely properties of Unbipentium
[2]
References [1] Sharma, M. (2005). "α-decay properties of superheavy elements Z=113-125 in the relativistic mean-field theory with vector self-coupling of ω meson". Physical Review C 71: 054310. doi: 10.1103/PhysRevC.71.054310 (http:/ / dx. doi. org/ 10. 1103/ PhysRevC. 71. 054310). [2] http:/ / www. apsidium. com/ elements/ 125. htm
Article Sources and Contributors
Article Sources and Contributors Unbipentium Source: http://en.wikipedia.org/w/index.php?oldid=306040099 Contributors: Bob Saint Clar, Clarityfiend, Drjezza, Femto, Headbomb, Isaac123456789, Kwamikagami, Nergaal, Qurqa, Shorelander, Stirling Newberry, Tallasse, Tony Sidaway, 14 anonymous edits
Image Sources, Licenses and Contributors Image:-TableImage.svg Source: http://en.wikipedia.org/w/index.php?title=File:-TableImage.svg License: GNU Free Documentation License Contributors: Bastique, Bibi Saint-Pol, Kwamikagami, Lantrix, Mdd4696, Popolon, Soeb, Tietew, 1 anonymous edits
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3
Unbihexium
1
Unbihexium 126
Ubp ← unbihexium → Ubs
↑
Ubh ↓
Ush periodic table - extended periodic table
General Name, Symbol, Number
unbihexium, Ubh, 126
Element category
Superactinides
Group, Period, Block
g6, 8, g
Appearance
unknown Standard atomic weight
−1
g·mol
Electron configuration
6
2
[Uuo] 5g 8s
Electrons per shell
2, 8, 18, 32, 38, 18, 8, 2 Physical properties
Phase
presumably solid Atomic properties
Oxidation states
presumably 3, 4, 6, 8 Miscellaneous Most-stable isotopes Main article: Isotopes of unbihexium
iso
NA
half-life
DM
DE (MeV)
DP
References
Unbihexium (pronounced /ˌuːnbaɪˈhɛksiəm/) is a hypothetical chemical element with atomic number 126 and symbol Ubh. It is of interest because it is in the hypothesized island of stability.
Unbihexium
History The first attempt to synthesize element 126 was performed in 1971 by Bimbot et al. using the hot fusion reaction:
A high energy alpha particle was observed and taken as possible evidence for the synthesis of element 126. Recent research suggests that this is highly unlikely as the sensitivity of experiments performed in 1971 would have been several orders of magnitude too low according to current understanding.
Stable unbihexium Calculations according to the Hartree-Fock-Bogoliubov Method using the non-relativistic Skyrme interaction have proposed Z=126 as a closed proton shell. In this region of the periodic table, N=184 and N=196 have been suggested as closed neutron shells. Therefore the isotopes of most interest are 310126 and 322126.
Predicted chemistry Element 126 is predicted to belong to a new block of valence g-electron atoms. The expected electron configuration is [Uuo]5g6 8s2 although there may be a smearing out of the energies of 5g and 7d orbitals. Recent calculations have indicated a stable monofluoride may exist, UbhF resulting from a bonding interaction between the 5g orbital on Ubh and the 2p orbital on fluorine.[1]
Unbihexium in popular culture An Action Comics story by John Byrne established the fictional Kryptonite as element 126 on the periodic table of the elements. The science fiction short story "Silence is Golden" [2] by American science fiction author Lou Antonelli is based on the supposed discovery of Element 126 in a Texas open-pit mine. The story, published by Revolution Science Fiction in August 2003, received an Honorable Mention in "The Year's Best Science Fiction, 21st annual collection" (St. Martin's Press, New York, N.Y. Gardner Dozois, ed. 2004). In the book "Travelers Rest", by Daniel Archangel, element 126 is used by humans to send carbon dioxide and water to aliens (Travelers) in space. Element 126 allows the spaceship, containing water and carbon dioxide, to travel into space without radiation harming the water and carbon dioxide.
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Unbihexium
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See also • Island of stability : Ununquadium – Unbinilium – → Unbihexium • Unbipentium – Unbiseptium • Period 8 element
External links • Likely properties of Unbihexium
[3]
References [1] Jacoby, Mitch (2006). " As-yet-unsynthesized superheavy atom should form a stable diatomic molecule with fluorine (http:/ / pubs. acs. org/ cen/ news/ 84/ i10/ 8410notw9. html)". Chemical & Engineering News 84 (10): 19. . Retrieved 2008-01-14. [2] http:/ / www. revolutionsf. com/ article. php?id=1959 [3] http:/ / www. apsidium. com/ elements/ 126. htm
Article Sources and Contributors
Article Sources and Contributors Unbihexium Source: http://en.wikipedia.org/w/index.php?oldid=303088335 Contributors: Abagofajay, Andros 1337, Antony-22, Anypodetos, Benjiboi, Biblbroks, BlueEarth, Bob Saint Clar, Calvero JP, CharlotteWebb, Complex01, Cosmium, DMacks, Dancanm, David Latapie, Dcljr, Deryck Chan, Drjezza, Długosz, Edgar181, Emk (ja), Eric119, Headbomb, INVERTED, Ilmari Karonen, Karlchwe, Kay Dekker, Kelovy, Kurykh, Kwamikagami, Materialscientist, Nergaal, Nimbulan, Odn2se, Oscarcwk, Poolkris, Pskykosys, RJHall, Reyk, Ruleke, Shanes, Sjlegg, Spug, Stirling Newberry, Talon Artaine, The Saad, Tony Sidaway, Van helsing, Warofdreams, Watch37264, 53 anonymous edits
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Unbioctium
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Unbioctium 128
Ubs ← unbioctium → Ube
↑
Ubo ↓
Uso periodic table - extended periodic table
General Name, Symbol, Number
unbioctium, Ubo, 128
Element category
Superactinides
Group, Period, Block
g8, 8, ?
Appearance
unknown - silvery or grey in color Standard atomic weight
−1
[340] g·mol Physical properties
Phase
presumably solid Miscellaneous Most-stable isotopes Main article: Isotopes of unbioctium
iso
NA
half-life
DM
DE (MeV)
DP
References
Unbioctium (pronounced /ˌuːnbaɪˈɒktiəm/) is the temporary name of a hypothetical undiscovered element on the periodic table that has the temporary symbol Ubo and atomic number 128.
Stability The element unbioctium would likely have isotopes that have a half life of only a few milliseconds, as this element is very close to an island of stability. An estimated five isotopes would then have half lives of over a half second while Ubo-340 would be the most stable.
Properties The physical properties of unbioctium, if it exists, are unknown, though this element may be expected to be highly reactive and oxidize quickly. It would presumably be a metal and solid. The color would probably be silvery white.
Unbioctium
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See also • Unbiseptium−Unbiennium
External links • Likely properties of unbioctium
[1]
References [1] http:/ / www. apsidium. com/ elements/ 128. htm
Article Sources and Contributors
Article Sources and Contributors Unbioctium Source: http://en.wikipedia.org/w/index.php?oldid=309175800 Contributors: Biblbroks, Bob Saint Clar, Choihei, Cosmium, DMacks, GraYoshi2x, Kurykh, Kwamikagami, Molinari, Nergaal, Rursus, SlayerK, Tony Sidaway, Watch37264, Weggie, 12 anonymous edits
Image Sources, Licenses and Contributors Image:-TableImage.svg Source: http://en.wikipedia.org/w/index.php?title=File:-TableImage.svg License: GNU Free Documentation License Contributors: Bastique, Bibi Saint-Pol, Kwamikagami, Lantrix, Mdd4696, Popolon, Soeb, Tietew, 1 anonymous edits
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Untriseptium
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Untriseptium 137
Uth ← untriseptium → Uto
↑
Uts ↓
Uos periodic table - extended periodic table
General Name, Symbol, Number
untriseptium, Uts, 137
Element category
Superactinides
Group, Period, Block
g17, 8, g
Appearance
unknown Standard atomic weight
−1
[364] u (supposition) g·mol
Electron configuration
18
[Uuo] 5g
Electrons per shell
1
8s
2, 8, 18, 32, 50, 18, 8, 1 Physical properties
Phase
presumably solid
Miscellaneous Most-stable isotopes Main article: Isotopes of untriseptium iso
NA
half-life
DM
DE (MeV)
DP
References
Untriseptium (pronounced /ˌʌntraɪˈsɛptiəm/) is a hypothetical chemical element which has not been observed to occur naturally, nor has it yet been synthesised. Due to drip instabilities, it is not known if this element is physically possible. Its atomic number is 137 and symbol is Uts. The name untriseptium is a temporary IUPAC systematic element name. The name Feynmanium (symbol Fy) is also informally used, because of Feynman's presentation of the speed of light problem described below.
Untriseptium
History The name untriseptium is an IUPAC systematic element name, the temporary name and symbol assigned to newly-synthesized and not-yet-synthesized chemical elements. A transuranic element receives a permanent name and symbol only after its synthesis has been confirmed. Transuranic elements (those beyond uranium) are, except for microscopic quantities and except for plutonium, always artificially produced, and usually end up being named for a scientist or the location of a laboratory that does work in atomic physics.
Significance Bohr model breakdown The Bohr model exhibits difficulty for atoms with atomic number greater than 137, for the speed of an electron in a 1s electron orbital, v, is given by:
where Z is the atomic number, and α is the fine structure constant, a measure of the strength of electromagnetic interactions.[1] Under this approximation, any element with an atomic number of greater than 137 would require 1s electrons to be traveling faster than c, the speed of light. Hence the non-relativistic Bohr model is clearly inaccurate when applied to such an element.
The Dirac equation The relativistic Dirac equation also has problems for Z > 137, for the ground state energy is
where m is the rest mass of the electron. For Z > 137, the wave function of the Dirac ground state is oscillatory, rather than bound, and there is no gap between the positive and negative energy spectra, as in the Klein paradox.[2] More accurate calculations including the effects of the finite size of the nucleus indicate that the binding energy first exceeds 2mc2 for Z > Zcr ≈ 137. For Z > Zcr, if the innermost orbital is not filled, the electric field of the nucleus will pull an electron out of the vacuum, resulting in the spontaneous emission of a positron.[3] More complete analysis involving relativity shows that the contradiction this particle poses may actually occur in the hypothetical 139-element (see the Wikipedia article unsolved problems in chemistry).
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Untriseptium
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See also • • • •
Dubnium Eka-Rutherfordium — Eka-Dubnium — Eka-Seaborgium Untrihexium – Untrioctium Atomic orbital, section Relativistic effects
External links • Feynman Online
[4]
References [1] See for example R. Eisberg and R. Resnick, Quantum Physics of Atoms, Molecules, Solids, Nuclei and Particles, Wiley (New York: 1985). [2] James D. Bjorken and Sidney D. Drell, Relativistic Quantum Mechanics, McGraw-Hill (New York:1964). [3] Walter Greiner and Stefan Schramm, Am. J. Phys. 76, 509 (2008), and references therein. [4] http:/ / www. fotuva. org/ online/ frameload. htm?/ online/ 137. htm
Article Sources and Contributors
Article Sources and Contributors Untriseptium Source: http://en.wikipedia.org/w/index.php?oldid=308834717 Contributors: ABCD, AWeishaupt, Army1987, BPinard, Bennetto, Biblbroks, BlueEarth, BlueMoonlet, Bob Saint Clar, Chunchulim, David Latapie, Eighteen and a half, Eric119, Giftlite, GregorB, HEL, Headbomb, INVERTED, Ixfd64, Jumping cheese, Kelovy, Kurykh, Kwamikagami, LukeSurl, NeilTarrant, Nick Y., NickBush24, Nightstallion, Oscarcwk, PJTraill, Poolkris, Rune.welsh, StradivariusTV, That Guy, From That Show!, Timc, Where, Yill577, 34 anonymous edits
Image Sources, Licenses and Contributors Image:-TableImage.svg Source: http://en.wikipedia.org/w/index.php?title=File:-TableImage.svg License: GNU Free Documentation License Contributors: Bastique, Bibi Saint-Pol, Kwamikagami, Lantrix, Mdd4696, Popolon, Soeb, Tietew, 1 anonymous edits
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