1. 1 CH5. Oxidation and Reduction

2. 2 History 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 MO = naturally occurring ores like ZnO, Fe 2 O 3, cuprates Separate into 2 reactions: (a) C(s) + O 2 (g)  CO 2 (g) (d) MO(s)  M (s or l) + ½ O 2 (g)  limit O 2

3. 3 Ellingham diagram Other possible reactions are: (b)C(s) + ½ O 2 (g)  CO(g) (c) CO(g) + ½ O 2 (g)  CO 2 (g) bronze = Cu/Sn alloy brass = Cu/Zn alloy

4. 4 Iron smelting

5. 5 Half-reactions 2H + (aqu) + 2e   H 2 (g)  G f   0 [H + ] = 1 H 2 pressure = 1 atm shorthand notation is H + /H 2 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. 6 Standard cell and potentials

7. 7 Half-reactions 2. Reverse rxn, reverse sign e  + A  A  E  = + 2V A   A + e  E  =  2V 3.Spontaneous rxns (  G neg) have positive potentials 4. Stoichiometry changes  G, not E 0 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 B 2   B + 2e  E  =  2.2V,  G = + 425 B 2  + 2A  2A  + B E  =  0.2V,  G = 425  2(190) = + 45

8. 8 Nernst equation 6. Nernst equation E = E   (0.059 / n) log Q Q = reaction quotient, for aA  bB + cC ; Q = [B] b [C] c / [A] a Ex: 2H 2 O  O 2 (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)  H 2 (g) E  = 0 V at pH=0 E = 0V  (0.059/2) log 1/[H + ] 2 = 0V  0.059(  pH) =  0.12V

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

10. 10 Kinetic factors Some redox reactions have slow kinetics, rates can be increased when overall E rxn > 0.6V (high overpotential exists) Converse statement – kinetically slow reactions may not occur at appreciable rates if E rxn < 0.6 V Examples of rapid reactions: 1. E rxn > 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  C 5 H 5 ) 2 ] +  [Fe(  5  C 5 H 5 ) 2 ] E  = 0.31V ferrocenium ferrocene

11. 11 Kinetic factors Examples of slow reactions: 1.E rxn < + 0.6V 2.Reactions that make/break strong bonds Ex. reactions with H 2, N 2, O 2 (water redox chemistry, N 2 fixation) 3.Reactions where n > 1 Ex: stability of MnO 4  in aqu acid MnO 4  / Mn 2+ E  = +1.51V at pH=0 4 ( 5e  + MnO 4  (aqu) + 8H + (aq)  Mn 2+ (aqu) + 4H 2 O ) + 1.51V 5 ( 2H 2 O  4e  + O 2 (g) + 4H + (aqu) ) - 1.23V 4MnO 4  (aqu) + 12H + (aqu)  4Mn 2+ (aqu) + 6H 2 O + 5O 2 (g) + 0.28V

12. 12 Kinetic factors 4.surface passivation Ex: Al anodization ~pH = 7 2Al(s) + 6OH  (aqu)  Al 2 O 3 (s) + 3H 2 O + 6e  E  ~ 1.7V ~ 1  m Al 2 O 3 passive surface forms during reaction and acts as a barrier to OH - and O 2 Ex: Si(m) in air forms a ~30nm SiO 2 native oxide passivation layer http://nano.boisestate.edu/research-areas/gate-oxide-studies/ Gate 1.0 nm SiO 2 on Si

13. 13 Combining half-rxns Combining red + red (or ox + ox) half-reactions: E  / V  G  / kJ/mol 1. e  + Mn 3+  Mn 2+ 1.5  148 2. e  + MnO 2 + 4H +  Mn 3+ + 2H 2 O 0.95  92 3. 2e  + MnO 2 + 4H +  Mn 2+ + 2H 2 O 1.23  240 E 3 = (n 1 E 1 + n 2 E 2 ) / n 3 = [(1)(1.5) + (1)(0.95)] / 2 = 1.23V Combining red + ox half-reactions: 1.e  + Mn 3+  Mn 2+ +1.5V 2.2H 2 O + Mn 3+  e  + MnO 2 + 4H +  0.95V 3.2H 2 O + 2Mn 3+  Mn 2+ + MnO 2 + 4H + +0.55V this disproportionation is spontaneous in acidic soln, but slow

14. 14 Latimer & Frost diagrams for Mn in acid HMnO 4 H 2 MnO 4 HMnO 3 MnO 2 Mn 3+ Mn 2+ Mn 1.51 2.091.23 1.69 0.90 1.28 2.9 0.95 1.5 -1.18

15. 15 Frost diagrams prop to -  G 

16. 16 Frost diagrams

17. 17 Frost diagram for N

18. 18 pH effect Oxoacids are better oxidants in acidic solution than in basic solution 10e  + 2HNO 3 + 10H +  N 2 + 6H 2 O E = 1.25V at pH=0 10e  + 2NO 3 - + 6H 2 O  N 2 + 12OH  E = 0.25V at pH=14 because they combine with H + to lose oxo or hydroxy ligands

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

20. 20 Pourbaix diagram for Fe e - + Fe 3+ → Fe 2+ E = +0.77 V e - + Fe(OH) 3 + 3H + → Fe 2+ + 3H 2 O E = E 0 - 3(0.059) pH e - + Fe(OH) 3 → Fe(OH) 2 + OH - E = E 0 - 0.059 pH

21. 21 Pourbaix diagram for Mn

22. 22 Example – Group 13