1. Knowing Nernst: Non-equilibrium copper redox chemistry

2. Objectives: (1)Calculate/measure stability of copper complexes (2)Use ligands to change stabilities of metal species HSAB concept: qualitative insights Redox potentials/Nernst eqn: quantitative insights

3. Chemical species studies CuCl 2 CuI Cu(NH 3 ) 4 2+ Cu(en) 2 2+ Cu(salen) n+ Charge vs oxidation state

4. Oxidation states Sum of oxidation states = ionic charge on species Assumes unequal sharing of electrons –more electronegative atom gets all of bond electrons

5. Oxidation states Sum of oxidation states = ionic charge on species Assumes unequal sharing of electrons –more electronegative atom gets all of bond electrons Examples: –MnO, MnO 2, KMnO 4 What differences are found between compounds with difference oxidation numbers? Atomic radius Reactivity (redox potential)

6. Disproportionation 2 Fe 4+ →Fe 3+ + Fe 5+ 2 H 2 O 2 → 2 H 2 O + O 2 2 Cu + →Cu 0 + Cu 2+ Reverse of process: comproportionation

7. Sample redox potential calculation CuCl 2 + ammonia -> Cu(NH 3 ) 4 2+ + chloride (1) Cu 2+ + Iˉ + eˉ  CuI0.86V (2)Cu 2+ + Clˉ + eˉ  CuCl0.54V (3)I 2 + 2eˉ  2Iˉ0.54V (4)Cu + (aq) + eˉ  Cu(s)0.52V (5)Cu 2+ (aq) + 2eˉ  Cu(s)0.37V (6)CuCl + eˉ  Cu(s) + Clˉ0.14V (7)Cu(NH 3 ) 4 2+ + 2eˉ  Cu(s) + 4NH 3 -0.12V (8) Cu 2+ (aq) + eˉ  Cu + (aq)-0.15V (9)CuI + eˉ  Cu(s) + Iˉ-0.19V (10)Cu(en) 2 2+ + 2eˉ  Cu + 2en-0.50V

8. Reduction:Cu 2+ (aq) + 2eˉ  Cu(s)E 0 = +0.37V(5) Oxidation:Cu(s) + 4NH 3  Cu(NH 3 ) 4 2 + + 2eˉE 0 = +0.12V(7*) Net:Cu 2+ (aq) + 4NH 3  Cu(NH 3 ) 4 2+ E 0 = +0.49V  G 0 = -nFE 0 n = mol e - F = 96,500 C / mol e - E 0 = standard reduction potential in V (1M conc, 1 atm pressure) 1 Joule = (1 Volt)(1 Coulomb)

9. Nernst Equation n = number of mol e - R = 8.3145 J/K-mol F = 96,500 C / mol e - E 0 = standard reduction potential in V (1M conc, 1 atm pressure) at 298 K

10. Hard vs. soft Describes the general bonding trends of chemical species (Lewis acids / Lewis bases) Hard acids prefer to bind to hard bases, while soft acids prefer to bind to soft bases

12. K stability = [AB] / [A][B] softerharder most stable complexes least stable complexes

13. Hard: low polarizability, primarily ionic bonding Soft: high polarizability, primarily covalent bonding

14. Lewis acids and bases Hard acids H +, Li +, Na +, K +, Rb +, Cs + Be 2+, Mg 2+, Ca 2+, Sr 2+, Ba 2+ BF 3, Al 3+, Si 4+, BCl 3, AlCl 3 Ti 4+, Cr 3+, Cr 2+, Mn 2+ Sc 3+, La 3+, Ce 4+, Gd 3+, Lu 3+, Th 4+, U 4+, Ti 4+, Zr 4+, Hf 4+, VO 4+, Cr 6+, Si 4+, Sn 4+ Borderline acids Fe 2+, Co 2+, Ni 2+, Cu 2+, Zn 2+ Rh 3+, Ir 3+, Ru 3+, Os 2+ R 3 C +, Sn 2+, Pb 2+ NO +, Sb 3+, Bi 3+ SO 2 Soft acids Tl +, Cu +, Ag +, Au +, Cd 2+ Hg 2+, Pd 2+, Pt 2+, M 0, RHg +, Hg 2 2+ BH 3 CH 2 HO +, RO + Hard bases F -, Cl - H 2 O, OH -, O 2- CH 3 COO -, ROH, RO -, R 2 O NO 3-, ClO 4- CO 3 2-, SO 4 2-, PO 4 3- NH 3, RNH 2 N 2 H 4 Borderline bases Br - NO 2-, N 3- SO 3 2- C 6 H 5 NH 2, pyridine N 2 Soft bases H -, I - H 2 S, HS -, S 2-, RSH, RS-, R 2 S SCN - (bound through S), CN -, RNC, CO R 3 P, C 2 H 4, C 6 H 6 (RO) 3 P

16. Experimental Details --Part G: watch out for oil drips and ethanol flames --do not throw away stir bars--recover them --dissolve all of the H 2 salen and Cusalen--no precipitates