Understanding Periodic Properties of Elements: Multielectron Atoms Explanation

1. Ch. 8 Periodic Properties of the Elements Multielectron Atoms • “Hydrogen-like” orbitals are used for all atoms • Energy levels are affected by other electrons • Coulomb’sLaw—electrostatic repulsion of like charges is proportional to the amount of charge, and inversely proportional to the distance between them (see text for eqn) • Shielding—screening of one electron from the nuclear charge by other electrons around the same atom • Penetration—probability of the electron to be close to the nucleus • Effective nuclear charge (Zeff)—the amount of nuclear charge an electron experiences after taking shielding into account • Degenerate—of equal energy

2. Order of Filling Subshells

3. Electron Spin and the Pauli Exclusion Principle • Electrons have intrinisic angular momentum -- “spin” -- ms • Possible values: ms = +1/2 and -1/2 (only two possible values) • Pauli Exclusion Principle: • No two electrons in an atom can have identical values of all 4 quantum numbers -- maximum of 2 electrons per orbital! • A single orbital can hold a “pair” of electrons with opposite “spins” • e.g. the 3rd shell (n = 3) can hold a maximum of 18 electrons: n = 3 l = 0 1 2 subshell 3s 3p 3d # orbitals 1 3 5 # electrons 2 6 10 = 18 total • A single electron in an orbital is called “unpaired” • Atoms with 1 or more unpaired electrons are paramagnetic, otherwise they are diamagnetic

4. Electronic Configurations • The Aufbau Principle -- Order of Filling Subshells • Atomic # = # of protons = # electrons (in neutral atom) • Add electrons to atomic orbitals, two per orbital, in the general order of increasing principle quantum number n, for example:

5. Hund’s Rule • Maximum number of unpaired electrons in orbitals of equal energy Orbital diagrams: C __ __ __ __ __ N __ __ __ __ __ O __ __ __ __ __ 1s 2s 2p 1s 2s 2p 1s 2s 2p

6. Relationship to Periodic Table e.g. complete electronic configuration of Ge (#32, group IV) Ge 1s22s22p63s23p64s23d104p2 or, Ge 1s22s22p63s23p63d104s24p2 (by values of n) • Short-hand notation-- show preceding inert gas config. • Ge [Ar]4s23d104p2 where [Ar] = 1s22s22p63s23p6

7. Valence Shell Configurations • valence shell-- largest value of n (e.g. for Ge, n = 4) plus any partially filled subshells Ge 4s24p2 (valence shell electron configuration) Ge __ __ __ __ (valence shell orbital diagram) Elements in same group have same valence shell e– configurations e.g. group V: N 2s22p3 P 3s23p3 As 4s24p3 Sb 5s25p3 Bi 6s26p3 4s 4p

8. Sample Questions • Write the complete electron configuration of gallium. Answer: • Write the short-hand electron configuration for zirconium. Answer: • Write the orbital diagram for the valence shell of tellurium. Answer:

9. Sample Question How many unpaired electrons does a ruthenium(II) ion, Ru2+, have? Show an appropriate orbital diagram to explain your answer. Is the atom paramagnetic or diamagnetic?

10. Variation of Atomic Properties Atomic Size (atomic radius, expressed in pm -- picometers) e.g. group 1 metals: e.g. some elements in 2nd period: (10–12 m!)

11. General Trend in Atomic Size Relative sizes of ions cations are smaller than parent atoms e.g. Na 186 pm 2s22p63s1 Na+ 95 pm 2s22p6 anions are larger than parent atoms e.g. Cl 99 pm 3s23p5 Cl– 181 pm 3s23p6

12. Ionization Energy I.E. = energy required to remove an electron from an atom or ion (always endothermic, positive values) e.g. Li(g) --> Li+(g) + e– I.E. = 520 kJ/mole Exceptions: special stability of filled subshells, and of half-filled subshells

13. Electron Affinity • E. A. = energy released when an electron is added to an atom or ion (usually exothermic, negative EA values) e.g. Cl(g) + e– --> Cl–(g) E. A. = -348 kJ/mol • The general trends in all these properties indicate that there is a special stability associated with filled-shell configurations. • Atoms tend to gain or lose an electron or two in order to achieve a stable “inert gas configuration” -- many important consequences of this in chemical bonding.

14. Types of Elements Metals: Shiny, malleable, ductile solids with high mp and bp Good electrical conductors Metal character increases to lower left of periodic table Nonmetals: Gases, liquids, or low-melting solids Non-conductors of electricity Diatomic elements: H2, O2, N2, F2, Cl2, Br2, I2 Metalloids: Intermediate properties, often semiconductors

15. Sample Questions Of the following atoms, circle the one with the highest electron affinity. K Cl P Br Na Write a balanced chemical equation that corresponds to the electron affinity of the element that you selected above.

16. Alkali Metals • They want to be +1! • Easily oxidized, low EA, low IE. • Density increases moving down the group. (mass rises faster than atomic radius) • Reactions • With halogens to form salts, e.g. 2 Na(s) + Cl2(g) 2 NaCl(s) • With water to make base + hydrogen, e.g. 2 K(s) + 2 H2O(l)  2 K+(aq) + 2 OH–(aq) + H2(g) • Reactions are more vigorous as you get lower in the group (why?) http://www.youtube.com/watch?v=9bAhCHedVB4&feature=relmfu http://www.youtube.com/watch?v=rtNaEFXOdAc&feature=relmfu

17. Halogens • They want to be –1! • Easily reduced, high EA, high IE. • Density increases moving down the group. (mass rises faster than atomic radius) • Reactions • With metals to form metal halides, e.g. 2 Al(s) + 3 Cl2(g) 2 AlCl3(s) • With hydrogen to form hydrogen halides (binary acids!), e.g. H2(g) + I2(s)  2 HI(g) • With other halogens to form interhalogen compounds, e.g. Br2(l) + F2(g)  2 BrF(g) • http://www.youtube.com/watch?v=F4IC_B9i4Sg

18. Noble Gases • Closed-shell electron configuration; very unreactive! • Used for lights, airtanks for divers, cryogens • Few reactions! Fluorides, oxides can be made under severe conditions. • Helium--helios (sun) • Krypton--kryptos (hidden) • Neon--neos (new) • Xeno--xenos (stranger)