CHAPTER 3: THE PERIODIC TABLE I.
II.
The Periodic Table A. Periods and Groups 1. Groups – columns in the periodic table, gives number of valence electrons 2. Period – rows in the periodic table, gives number of occupied energy levels 3. Gab between group 2 and group 3 is filled by transition elements from the 4th period onwards Physical Properties A. Nuclear Charge – given by the atomic number, and increases by one between successive elements on the table as a pr oton is added to the nucleus B. Effective Nuclear Charge – charge felt by the valence electrons after taking into account the number of shielding electrons that surround the nucleus. 1. Effective nuclear charge increases with nuclear charge as there is no change in the number of inner electrons 2. Effective nuclear charge experienced by an atom’s outer electrons increases with the group number of the element. It increases across a period, but remains the same down the group. C. Shielding Electrons – the electrons in energy levels between the nucleus and valence electrons. They are known as shielding electrons because they shield the valence electrons from the full force of attraction exerted by the positive charge in the nucleus. D. Atomic Radius –half the distance between neighboring nuclei, and can also be considered as the distance from the nucleus to the other most electrons of the Bohr atom. 1. Atomic radii increase down a group as the number of occupied electron electron shells (given by period number) increases 2. Atomic radii decrease across a period E. Ionic Radius (self-explanatory) 1. Positive ions are smaller than their parent atoms as the formation of positive ions involves the loss of other shells. Cation radii is always smaller than radius of the atoms from which they are derived as when metals are converted to their cations, they always do so by losing the electrons in their highest energy levels. Since the ion has less electron that the atom from which it was derived, there is less mutual repulsion among electrons, so electron orbitals shrink. 2. Negative ions are larger than parent atoms, as the formation of negative ions involves addition of electrons into outer shells. Increased electron repulsion between electrons in outer shell causes the electrons to move further apart, and increase the radius of the outer shell 3. Ionic radii increase down a group as the number of electron shells increase 4. Ionic radii decreases across a period F. Ionization Energy 1. Ionization energy is the amount of energy required to remove an electron from an atom to form a positive ion 2. Ionization energy increases across a period – increase in effective nuclear charge causes an increase in attraction between outer electrons and nucleus makes electrons more difficult.
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3. Decreases down a group, as electron removed is an electron shell furthest from the nucleus. Nuclear charge increase, but effective nuclear charge is about the same, due to electron shielding, so increased distance reduces attraction G. Electronegativity – strength with which an atom pulls on the electrons it shares in a covalent bond 1. Electronegativity increases from left to right across a period 2. Increase from top to bottom in a group 3. Positively charged nucleus pulls on the other electrons, so the smaller an atom, the higher the electronegativity 4. Noble gases do not obey trend due to the stability they gain from having full other shells H. Melting Points 1. Depends on type of bonding and structure 2. Melting points decreases down group 1, they have metallic structures head by attractive forces between delocalized other electrons and positively charged ions. Attraction decreases with distance 3. Melting points increases down group 7 Chemical Properties A. Group 0: Noble Gases 1. Colorless, monoatomic, unreactive 2. Do not form positive ions, as they have highest ionization energies. Do not form negative ions, as extra electrons added to empty outer shell would experience negligible effective nuclear force as protons shield equal number of inner electrons. They all have stable octet. 3. Other elements gain or lose electrons as to achieve electron arrangement of nearest noble gas B. Group 1: Alkali Metals 1. Too reactive to be found in nature 2. Good conductors of electricity, low density, grey shiny surfaces when freshly cut, very reactive, forms ionic compounds with non-metals 3. Reaction with Water i. Reacts with water to produce hydrogen and the metal hydroxide ii. Reaction becomes more vigorous as we descend the group C. Group 7: Halogens 1. Exists as diatomic molecules 2. Nuclei have high effective charge and exert strong pull on any electron from other atoms. This electron can occupy outer energy level of the halogen atom and complete a stable octet 3. Attraction is greatest for smallest atom fluorine, which is most reactive among halogens 4. Reactivity decreases down the group, as atomic radius increases, and attraction for other electrons decreases 5. Reactions with Alkali Metals i. Halogens react with alkali metals to form ionic halides ii. Halogens form insoluble salts with silver. Adding solution containing halide to a solution containing silver ions produces a precipitate which is useful in identifying the halide Trends Across Period 3
V.
First row d-block Elements
