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’s Law—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 (Z eff )—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” -- m s –Possible values: m s = +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 = 012 subshell3s3p3d # orbitals135 # electrons2610 = 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: #AtomConfiguration 1H1s 1 2He1s 2 3Li1s 2 2s 1 4Be1s 2 2s 2 5B1s 2 2s 2 2p 1 6C1s 2 2s 2 2p 2 7N1s 2 2s 2 2p 3 8O1s 2 2s 2 2p 4 9F1s 2 2s 2 2p 5 10Ne1s 2 2s22p 6 11Na1s 2 2s 2 2p 6 3s 1

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

6. Relationship to Periodic Table e.g. complete electronic configuration of Ge (#32, group IV) Ge 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 2 or, Ge 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 2 (by values of n) Short-hand notation -- show preceding inert gas config. –Ge [Ar]4s 2 3d 10 4p 2 where [Ar] = 1s 2 2s 2 2p 6 3s 2 3p 6

7. Valence Shell Configurations valence shell -- largest value of n (e.g. for Ge, n = 4) plus any partially filled subshells Ge 4s 2 4p 2 (valence shell electron configuration) Ge __ __ __ __ (valence shell orbital diagram) Elements in same group have same valence shell e – configurations e.g. group V: N 2s 2 2p 3 P3s 2 3p 3 As4s 2 4p 3 Sb5s 2 5p 3 Bi6s 2 6p 3 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, Ru 2+, 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: AtomRadius in pmValence Shell Li1522s 1 Na1863s 1 K2274s 1 Cs2485s 1 AtomBCNOF radius8877706664 e – config2p 1 2p 2 2p 3 2p 4 2p 5 (10 –12 m!)

11. General Trend in Atomic Size Relative sizes of ions cations are smaller than parent atoms e.g. Na 186 pm 2s 2 2p 6 3s 1 Na + 95 pm 2s 2 2p 6 anions are larger than parent atoms e.g. Cl 99 pm 3s 2 3p 5 Cl – 181 pm 3s 2 3p 6

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 Metalloids: Intermediate properties, often semiconductors Diatomic elements: H 2, O 2, N 2, F 2, Cl 2, Br 2, I 2

15. Sample Questions Of the following atoms, circle the one with the highest electron affinity. KClPBrNa 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) + Cl 2(g)  2 NaCl (s) –With water to make base + hydrogen, e.g. 2 K (s) + 2 H 2 O (l)  2 K + (aq) + 2 OH – (aq) + H 2(g) Reactions are more vigorous as you get lower in the group (why?) http://www.youtube.com/watch?v=9bAhCHedVB4&feature=rel mfu http://www.youtube.com/watch?v=rtNaEFXOdAc&feature=rel mfu

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 Cl 2(g)  2 AlCl 3(s) –With hydrogen to form hydrogen halides (binary acids!), e.g. H 2(g) + I 2(s)  2 HI (g) –With other halogens to form interhalogen compounds, e.g. Br 2(l) + F 2(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)