mega (M)106 kilo (k) 10³ hecto (h) 10² deca (da) 101 1.0 deci (d) 10-1 centi (c) 10-2 milli (m) 10-3 micro (µ) 10-6 nano (n) 10-9 pico (p) 10-12 SI Basics: ), vol(L), L-intensity(cd) Mass(kg) length(m) time(s) temp(k) amount(mol) e-current(A:ampere ), energy(J ), J=(kgm2)/s2=Nm E= (hc)/
h= Planck : 6.626 E-34Js
c: Speed of Light= 2.998 E8m/s 1L=(1dm)3
D=m(g)/v(cm3)(g/cm3)
1mol=6.022E23
wl
1 angstrom Å=1.0E-10 m 3
i.e. 1L= (10cm) 1000cm
(1mol amu= 6.022E23 amu = 1g) TC = (TF - 32°F) (5°C/9°F)
(m)=
3
Å=0.1nm
1cm3= 1mL
1amu= 1g/mol (atomic mass u)
k = °C + 273.15
°F = °C + 32°
Measurement- quantitative amount observed (composed of # & a scale or unit)) Accuracy- agreement of measurement w/ true value
measurements Precision- degree of agreement btwn repeated measurements Random err- varying magnitude, measurement of being h or l Systematic err- occurs in same direction & amount each time (ruler) Zero’s: Leading: 0.0252sf Exact #’s: ∞ sig
Captive:1.08-3sf)
:101sf &10.2sf Trailing :10
figs (i.e: counting #’s, 1in=2.54cm, & 2πr )
1cm3 = 1mL, 1dm3 = 1L, 1 Angstrom (Å)= 1.0 × 10-10 m 1mol= 6.022 E23 units, (1 mol amu= 6.022 E23 amu = 1 g), so 1amu = 1g/mol ° C + 273.15 = k, °F= °C + 32°, (TF - 32°F) 5° C/9°F=TC, D=m/v (g/cm3) Mixtures: Homogenous (vis. indistinguish) & hetero (vis. distinguish) Chemistry: how 1 substance changes to another
Science: understand nature & explain phenomena of physical world Scientific method: process at center of scientific inquiry:
1. Observe: collect data 2. Predict: hypoth 3. Experiment: test hypoth Observations: Qualitative no #, Quantitative- involves # & a unit Hypothesis: possible explanation for an observation, educated guess Theory/model: set of tested hypoth explanation of natural phenomena Law- summary of observed; same observ applies to diff systems Pure substance: constant composition. can be isolated by distil, filter, & chrom Compound: thing w/ constant comp. can be chem broken into elements i.e. electrolysis of H2O H & O Physical change: change in form of substance but not in chemical composition Chemical change: substance new substance(s) w/ diff properties & composition Law of conservation of mass - total mass of materials is not affected by a chemical change
R=PV/T for gas constant (R), pressure (P), volume (V), & temperature (T). law of definite prop- A given compound always contains exactly t he same proportion of elements by mass. law of multiple proportions: when 2 elements form a series of compounds, ratios of masses of 2nd element that combine w/ 1 gram of 1st element can always be reduced to small whole #’s Dalton’s Atomic Theory:
1. Each element is made of atoms. 2. The atoms of an element= identical; atoms of diff elements= diff in fundamental way(s) 3. Chem cmpnds form when atoms of diff elements combine w/ eachother. A given cmpnd alwys has same rel. #’s & types of atoms. 4. Chem reaction involves reorganizing atoms, changing way bound together). atoms are not changed in chemical reaction.
Avogadro’s hypothesis - at same temp & pressure, equal vol of diff gases contain same # of particles. 2H + 1O 2H20 charge-to-mass ratio: of an electron: e/m=-1.76 E8C/g [e= charge on electron (C) & m=electron mass(g)] 1)
As you move down a group, atomic radius increases. The number of energy levels increases as you move down a group as the number of electrons increases. Each subsequent energy level is further from the nucleus than the last. Therefore, atomic radius increases as the group & energy levels increase.
2) As you move across a period, atomic radius decreases. As you go across a period, electrons are added to the same energy level. At the same time, protons are being added to the nucleus. The concentration of more protons in the nucleus creates a " higher effective nuclear charge .
1) Anions (negative ions) are larger than their respective atoms. Electron-electron repulsion forces them to spread further apart. e- outnumber p+; protons cannot pull extra electrons as tightly toward nucleus. 2) Cations (positive ions) are smaller than their respective atoms. There is less electron-electron repulsion, so they can come closer together. Protons outnumber electrons; the protons can pull the fewer electrons toward the nucleus more tightly. First Ionization Energy: The energy required to remove the outermost (highest energy) electron from a neutral atom in its ground state. As you move down a group, first ionization energy decreases. Electrons are further from the nucleus and thus easier to remove the outermost one. " SHIELDING" - Inner electrons at lower energy levels essentially block the protons' force of attraction toward the nucleus. It therefore becomes easier to remove the outer electron As you move across a period, first ionization energy increases. As you move across a period, the atomic radius decreases, that is, the atom is smaller. The outer electrons are closer to the nucleus and more strongly attracted to the center. Therefore, it becomes more difficult to remove the outermost electron. Exceptions to First Ionization Energy Trends:
1) Xs2 > Xp1 e.g. 4Be > 5B The energy of an electron in an Xp orbital is greater than the energy of an electron in its respective Xs orbital. Therefore, it requires less energy to remove the first electron in a p orbital than it is to remove one from a filled s orbital. 2) Xp3 > Xp4 e.g. 7N > 8O After the separate degenerate orbitals have been filled with single electrons, the fourth electron must be paired. The electron-electron repulsion makes it easier to remove the outermost, paired electron. • Hund's Rule)
Second and Higher Ionization Energies
Definition: Second Ionization Energy is the energy required to remove a second outermost electron from a ground state atom. Subsequent ionization energies increase greatly once an ion has reached the state like that of a noble gas. In other words, it becomes extremely difficult to remove an electron from an atom once it loses enough electrons to lose an entire energy level so that its valence shell is filled. Ionization Energies (kJ/mol)
Element
1st IE
2nd IE
3rd IE
4th IE
Na
495.8
4562.4
6912
9543
Mg
737.7
1450.6
7732.6
10,540
Al
577.6
1816.6
2744.7
11,577
Electron Affinity- The energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion. 1)
As you move down a group, electron affinity decreases. 2) As you move across a period, electron affinity increases.
Exceptions: •Among nonmetals, however, the elements in the first period have lower electron affinities than the elements below them in their respective groups. •Elements w/ electron configurations of Xs2, Xp3, & Xp6 have electron affinities less than zero because they are unusually stable. In other words instead of energy being given off, these elements actually require an input of energy in order to gain electrons. i.e. Be, N, Ne WHY??? B/C Electron affinities are all much smaller than ionization energies. 1. Xs2 < 0: Stable, diamagnetic atom w/ no unpaired electrons. 2. Xp3 < 0: Stable atom w/ 3 unpaired p-orbital electrons each occupying its own subshell. 4. Xp6 < 0: Stable atom w/ filled valence (outermost) shell. mass of the electron = 9.11 E-31kg nuclear atom — an atom w/ a dense center of + charge (the nucleus) w/ electrons moving around the nucleus at a distance that is large relative to nuclear radius. Isotopes- atoms with the same # of protons but different #’ s of neutrons. Atomic number Z (# of protons) is written as subscript, & mass number A (total # of protons & neutrons) is written as a superscript.
Share electrons- covalent bond Ion: is an atom or group of atoms that has a net + or - charge. nonmetals appear in the upper-right corner of the table alkali metals, members of Group 1A — lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr) — are very active elements that readily form ions with a 1 charge when they react with nonmetals. alkaline earth metals- The members of Group 2A — beryllium (Be), magnesium (Mg), calcium ( Ca), strontium (Sr), barium (Ba), and radium (Ra. They all form ions with a 2 charge when they react with nonmetals. Halogens- members of Group 7A — fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At) — all form diatomic molecules. Fluorine, chlorine, bromine, and iodine all react with metals to form salts containing ions with a 1- charge (F-, Cl-, Br-, and I-).
The members of Group 8A — helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn) — are known as the noble gases. They all exist under normal conditions as monatomic (single-atom) gases and have little chemical reactivity. The SI def of mole is the amount of a substance that contains as many entities as there are in exactly 12 g of carbon-12.
Avogadro’s number is 6.022 E23. 1 mol of anything- 6.022 E 23 units of that substance.
