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Understanding Non-Equilibrium Copper Redox Chemistry: Nernst Equation Explained

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nigel-davis
17 diapositives · 299 Views · Oct 26, 2014

Explore the Nernst equation and its application to the study of copper redox chemistry, focusing on stability calculations, ligand effects, and oxidation states. Gain qualitative and quantitative insights into chemical reactions.

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Understanding Non-Equilibrium Copper Redox Chemistry: Nernst Equation Explained
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  1. Handouts • Syllabus • Safety (detailed) • Web page is up and running

  2. Knowing Nernst:Non-equilibrium copper redox chemistry

  3. Safety / Waste disposal • There will only be one liquid waste disposal bottle (kept in the hood by the lab entrance), since all reactions are carried out in aqueous solution. There should be a separate waste bottle for solid waste.

  4. 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

  5. Chemical species studies • CuCl2 • CuI • Cu(NH3)2+ • Cu(en)22+ • Cu(salen)n+ • Charge vs oxidation state

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

  7. 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, Mn2O3, Mn3O4, MnO2, Mn5O8, KMnO4 • What differences are found between compounds with difference oxidation numbers? Atomic radius Reactivity (redox potential)

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

  9. 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

  10. 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)=2+ E0 = +0.49V

  11. 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. most stable complexes Kstability = [AB] / [A][B] Monotonic variation in stability least stable complexes Only two possible trends harder softer

  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+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

  15. Topics: • Nernst equation (Electrochemistry problems)

  16. Hard/soft references • R.G. Pearson, Inorg. Chem., 27, p734 (1988). • R.G. Pearson, JACS, 85, p3533 (1963) • R.G. Pearson, J. Chem. Ed., 45, p581 AND p643 (1968) • R.G. Pearson, J. Chem. Ed. 64 (7): 561-567 JUL 1987 [471 cites] • Hard and soft acids and bases, Ralph G. Pearson, editor. (1973) 480pp • Hard and soft acids and bases {principle} in organic chemistry, T. L. Ho (1977) 209pp