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Knowing Nernst: Non-equilibrium copper redox chemistry

Knowing Nernst: Non-equilibrium copper redox chemistry. Knowing Nernst: Non-equilibrium copper redox chemistry. Objectives: Calculate/measure stability of copper complexes Use ligands to change stabilities of metal species. HSAB concept: qualitative insights

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Knowing Nernst: Non-equilibrium copper redox chemistry

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  1. Knowing Nernst:Non-equilibrium copper redox chemistry

  2. Knowing Nernst:Non-equilibrium copper redox chemistry Objectives: Calculate/measure stability of copper complexes Use ligands to change stabilities of metal species HSAB concept: qualitative insights Redox potentials/Nernst eqn: quantitative insights

  3. Chemical species studies • CuCl2 • CuI • Cu(NH3)42+ • Cu(en)22+ • 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, MnO2, KMnO4 • What differences are found between compounds with difference oxidation numbers? Atomic radius Reactivity (redox potential)

  6. Disproportionation • 2 Fe4+→ Fe3+ + Fe5+ • 2 H2O2 → 2 H2O + O2 • 2 Cu+ → Cu0 + Cu2+ • Reverse of process: comproportionation

  7. Sample redox potential calculation CuCl2 + ammonia -> Cu(NH3)42+ + chloride (1) Cu2+ + Iˉ + eˉ  CuI 0.86V (2) Cu2+ + Clˉ + eˉ  CuCl 0.54V (3) I2 + 2eˉ  2Iˉ 0.54V (4) Cu+ (aq) + eˉ  Cu(s) 0.52V (5) Cu2+(aq) + 2eˉ  Cu(s) 0.37V (6) CuCl + eˉ  Cu(s) + Clˉ 0.14V (7) Cu(NH3)42+ + 2eˉ  Cu(s) + 4NH3 -0.12V (8) Cu2+(aq) + eˉ  Cu+ (aq) -0.15V (9) CuI + eˉ  Cu(s) + Iˉ -0.19V (10) Cu(en)22+ + 2eˉ  Cu + 2en -0.50V

  8. Reduction: Cu2+(aq) + 2eˉ  Cu(s) E0 = +0.37V (5) Oxidation: Cu(s) + 4NH3 Cu(NH3)42+ + 2eˉ E0 = +0.12V (7*) Net: Cu2+(aq) + 4NH3 Cu(NH3)42+ E0 = +0.49V • DG0 = -nFE0 • n = mol e- • F = 96,500 C / mol e- • E0 = standard reduction potential in V (1M conc, 1 atm pressure) • 1 Joule = (1 Volt)(1 Coulomb)

  9. Nernst Equation at 298 K n = number of mol e- R = 8.3145 J/K-mol F = 96,500 C / mol e- E0 = standard reduction potential in V (1M conc, 1 atm pressure)

  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

  11. most stable complexes Kstability = [AB] / [A][B] least stable complexes harder softer

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

  13. Lewis acids and bases • Hard acids H+, Li+, Na+, K+ , Rb+, Cs+Be2+, Mg2+, Ca2+ , Sr2+, Ba2+BF3, Al 3+, Si 4+, BCl3 , AlCl3Ti4+, Cr3+, Cr2+, Mn2+Sc3+, La3+, Ce4+, Gd3+, Lu3+, Th4+, U4+, Ti4+, Zr4+, Hf4+, VO4+, Cr6+,  Si4+, Sn4+ • Borderline acids Fe2+, Co2+, Ni2+ , Cu2+, Zn2+Rh3+, Ir3+, Ru3+, Os2+R3C+ , Sn2+, Pb2+NO+, Sb3+, Bi3+SO2 • Soft acids Tl+, Cu+, Ag+, Au+, Cd2+Hg2+, Pd2+, Pt2+, M0, RHg+, Hg22+BH3CH2HO+, RO+ • Borderline bases Br-NO2-, N3- SO32-C6H5NH2, pyridine N2 • Soft bases H-, I-H2S, HS-, S2- , RSH, RS-, R2SSCN- (bound through S), CN-, RNC, CO R3P, C2H4, C6H6(RO)3P  • Hard bases F-, Cl-H2O, OH-, O2-CH3COO- , ROH, RO-, R2ONO3-, ClO4-CO32-, SO42- , PO43-NH3, RNH2N2H4

  14. Experimental Details --Part G: watch out for oil drips and ethanol flames --do not throw away stir bars--recover them --dissolve all of the H2salen and Cusalen--no precipitates

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