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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) + 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
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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
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Half-reactions2H+ (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
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Half-reactions2. 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 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
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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: 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
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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
Stability field for water
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Kinetic factorsSome 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(5C5H5)2]+ [Fe(5C5H5)2] E = 0.31V
ferrocenium ferrocene
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Kinetic factors
Examples of slow reactions:
1. Erxn < + 0.6V
2. Reactions that make/break strong bondsEx. reactions with H2, N2, O2 (water redox chemistry, N2 fixation)
3. Reactions where n > 1
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
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Kinetic factors4. 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
http://nano.boisestate.edu/research-areas/gate-oxide-studies/
Gate 1.0 nm SiO2 on Si
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Combining half-rxnsCombining 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
3. 2H2O + 2Mn3+ Mn2+ + MnO2 + 4H+ +0.55V
this disproportionation is spontaneous in acidic soln, but slow
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Latimer & Frost diagrams for Mn in acid
HMnO4 H2MnO4 HMnO3 MnO2 Mn3+ Mn2+ Mn
1.51
2.09 1.23
1.69
0.90 1.28 2.9 0.95 1.5 -1.18
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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
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Ligand effectsNote 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
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)
CN poisoning inhibits cytochrome oxidase in mitochondria
Zn(m)
Zn(CN)2(s) + Au(s)
KOH[Zn(OH)4]2(aqu) + Au(s)
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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
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Pourbaix diagram for Mn