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CH5. Oxidation and Reduction

CH5. Oxidation and Reduction. Redox chemistry involves changes in elemental oxidation states during reaction Historically – first man-made redox reactions might be forming metals 2 MO(s) + C(s)  2 M (s or l) + CO 2 (g) smelting

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CH5. Oxidation and Reduction

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  1. CH5. Oxidation and Reduction

  2. Redox chemistry involves changes in elemental oxidation states during reaction Historically – first man-made redox reactions might be forming metals 2 MO(s) + C(s)  2 M (s or l) + CO2(g) smelting MO = naturally occurring ores like ZnO, Fe2O3, cuprates Separate into 2 reactions: (a) C(s) + O2(g)  CO2(g) (d) MO(s)  M (s or l) + ½ O2(g)  limit O2 History

  3. Ellingham diagram Other possible reactions are: (b) C(s) + ½ O2(g)  CO(g) (c) CO(g) + ½ O2(g)  CO2(g) bronze = Cu/Sn alloy brass = Cu/Zn alloy

  4. Iron smelting

  5. Half-reactions • 2H+ (aqu) + 2e H2(g) Gf 0 • [H+] = 1 H2 pressure = 1 atm • shorthand notation is H+/H2 redox couple • 1. G = nFE • n = number of e transferred F = Faraday’s constant = 96480 C / mol E = std. potential for a rxn or half-rxn • E givesG and v.v. (thermodynamic data can be used to calc E) • note: 1 kJ = 1000 CV, so 1 eV  100 kJ/mol nEG

  6. Standard cell and potentials

  7. 2. Reverse rxn, reverse sign e + A  A E = + 2V • A A + e E =  2V • Spontaneous rxns (G neg) have positive potentials 4. Stoichiometry changes G, not E0 e + A  A E = + 2V, G = -190 kJ/mol 2e + 2A  2A E = + 2V, G = -380 5. Adding oxidation to reduction half-reactions 2 (e + A  A) E = +2V, G = 190 kJ/mol • B2  B + 2e E = 2.2V, G = + 425 • B2 + 2A  2A + B E = 0.2V, G = 425  2(190) = + 45 Half-reactions

  8. Nernst equation • . Nernst equation E = E (0.059 / n) log Q Q = reaction quotient, for aA  bB + cC ; Q = [B]b [C]c / [A]a • Ex: 2H2O  O2(g) + 4H+(aqu) + 4e E = 1.23V at pH=0 • Ex: What is the half-reaction potential to oxidize water at pH = 2? • E = E (0.059/4)log [H+]4 = 1.23V + 0.059(ΔpH) = -1.23V + 0.12V = -1.11V • Ex: What is the water reduction potential at pH = 2? • 2e + 2H+(aqu)  H2(g) E = 0 V at pH=0 • E = 0V  (0.059/2) log 1/[H+]2 = 0V  0.059(pH) = 0.12V

  9. Stability field for water Note that E(O2/H2O)  E(H2O/H2) = 1.23V (pH independent) E (V) 2e + 2H+(aqu)  H2(g) 0.00 - 0.059pH H2O  ½O2(g) + 2H+ + 2e -1.23 + 0.059pH H2O  H2(g) + ½ O2(g) -1.23V

  10. Kinetic factors Some redox reactions have slow kinetics, rates can be increased when overall Erxn > 0.6V (high overpotential exists) Converse statement – kinetically slow reactions may not occur at appreciable rates if Erxn < 0.6 V Examples of rapid reactions: 1. Erxn > 0.6V 2. outer-sphere mechanisms reaction does not make/break strong bonds or change coordination geometry Ex: e + [Fe(CN)6]3(aqu)  [Fe(CN)6]4(aqu) E = 0.38V hexacyanoferrate(III) hexacyanoferrate(II) ferricyanate ferrocyanate Ex: e + [Fe(5C5H5)2]+ [Fe(5C5H5)2]E = 0.31V ferrocenium ferrocene

  11. Ex: stability of MnO4 in aqu acid MnO4 / Mn2+ E = +1.51V at pH=0 4 ( 5e+ MnO4(aqu) + 8H+(aq)  Mn2+(aqu) + 4H2O ) + 1.51V 5 ( 2H2O  4e + O2(g) + 4H+(aqu) ) - 1.23V 4MnO4(aqu) + 12H+(aqu)  4Mn2+(aqu) + 6H2O + 5O2(g) + 0.28V Kinetic factors Examples of slow reactions: 1. Erxn < + 0.6V 2. Reactions that make/break strong bonds Ex. reactions with H2, N2, O2 (water redox chemistry, N2 fixation) Reactions where n > 1

  12. Kinetic factors surface passivation Ex: Al anodization ~pH = 7 2Al(s) + 6OH(aqu)  Al2O3(s) + 3H2O + 6e E ~ 1.7V ~ 1 m Al2O3 passive surface forms during reaction and acts as a barrier to OH- and O2 Ex: Si(m) in air forms a ~30nm SiO2 native oxide passivation layer Gate 1.0 nm SiO2 on Si http://nano.boisestate.edu/research-areas/gate-oxide-studies/

  13. Combining half-rxns • Combining red + red (or ox + ox) half-reactions: • E / VG / kJ/mol • 1. e + Mn3+ Mn2+ 1.5 148 • 2. e + MnO2 + 4H+ Mn3+ + 2H2O 0.95 92 • 3. 2e + MnO2 + 4H+  Mn2+ + 2H2O 1.23 240 • E3 = (n1E1 + n2 E2) / n3 = [(1)(1.5) + (1)(0.95)] / 2 = 1.23V • Combining red + ox half-reactions: • 1. e + Mn3+ Mn2+ +1.5V • 2. 2H2O + Mn3+ e + MnO2 + 4H+0.95V • 2H2O + 2Mn3+ Mn2+ + MnO2 + 4H+ +0.55V • this disproportionation is spontaneous in acidic soln, but slow

  14. 1.51 0.90 1.28 2.9 0.95 1.5 -1.18 HMnO4 H2MnO4 HMnO3 MnO2 Mn3+ Mn2+ Mn 2.09 1.23 1.69 Latimer & Frost diagrams for Mn in acid

  15. Frost diagrams prop to -G

  16. Frost diagrams

  17. Frost diagram for N

  18. pH effect Oxoacids are better oxidants in acidic solution than in basic solution 10e + 2HNO3 + 10H+ N2 + 6H2O E = 1.25V at pH=0 10e + 2NO3- + 6H2O  N2 + 12OH E = 0.25V at pH=14 because they combine with H+ to lose oxo or hydroxy ligands

  19. O2 + 4H+ + 4e 2H2O +1.23 2CN + Au  [Au(CN)2] + e0.60 O2 + 4H+ + 8CN + 4Au  4[Au(CN)2] + 2H2O E = +0.63 (pH=0) Ligand effects Note that e + Fe3+(aqu)  Fe2+(aqu) E = +0.77V But e + [Fe(CN)6]3(aqu)  [Fe(CN)6]4(aqu) E = +0.38V => cyano ligand stabilizes Fe3+ more than OH2 +1.80V +0.80 AgO  Ag+  Ag(m) pH=0 +0.60 +0.34 AgO  Ag2O  Ag(m) pH=14 +1.69 Au+  Au(m) pH=0 +0.60 [Au(CN)2]  Au(m) pH=0 Zn(m) Zn(CN)2(s) + Au(s) CN poisoning  inhibits cytochrome oxidase in mitochondria KOH [Zn(OH)4]2(aqu) + Au(s)

  20. Pourbaix diagram for Fe e- + Fe3+ → Fe2+ E = +0.77 V e- + Fe(OH)3 + 3H+ → Fe2+ + 3H2O E = E0 - 3(0.059) pH e- + Fe(OH)3→ Fe(OH)2 + OH- E = E0 - 0.059 pH

  21. Pourbaix diagram for Mn

  22. Example – Group 13

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