CH5. Oxidation and Reduction
1
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)
limit O2
smelting
MO = naturally occurring ores like ZnO, Fe2O3, cuprates
Separate into 2 reactions:
+ O2(g)  CO2(g)
(a)
C(s)
(d)
MO(s)  M (s or l) + ½ O2(g)
2
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
3
Iron smelting
4
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
5
Standard cell and potentials
6
Half-reactions
e + A  A
A  A + e
2.
Reverse rxn, reverse sign
E = + 2V
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
7
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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Stability field for water
Note that E(O2/H2O)  E(H2O/H2) = 1.23V (pH independent)
E (V)
2e + 2H+(aqu)  H2(g)
H2O  ½O2(g) + 2H+ + 2e
H2O  H2(g) + ½ O2(g)
0.00 - 0.059pH
-1.23 + 0.059pH
-1.23V
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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)
hexacyanoferrate(III)
ferricyanate
Ex:
e +
hexacyanoferrate(II)
ferrocyanate
[Fe(5C5H5)2]+ 
ferrocenium
E = 0.38V
[Fe(5C5H5)2]
E = 0.31V
ferrocene
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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)
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 factors
4.
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/
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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+
2.
2H2O + Mn3+  e + MnO2 + 4H+
0.95V
3.
2H2O + 2Mn3+  Mn2+ + MnO2 + 4H+
+0.55V
+1.5V
this disproportionation is spontaneous in acidic soln, but slow
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Latimer & Frost diagrams for Mn in acid
0.90
HMnO4
1.51
2.9
1.28
H2MnO4
HMnO3
2.09
0.95
MnO2
Mn3+
1.5
Mn2+
-1.18
Mn
1.23
1.69
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Frost diagrams
prop to
-G
15
Frost diagrams
16
Frost diagram for N
17
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
18
Ligand effects

Note that
e +
Fe3+(aqu)
But
e +
[Fe(CN)6]3(aqu)  [Fe(CN)6]4(aqu)
Fe2+(aqu)
E = +0.77V
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(m)
Au+
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)
Zn(m)
Zn(CN)2(s) + Au(s)
poisoning  inhibits
cytochrome oxidase in mitochondria
CN
KOH
[Zn(OH)4]2(aqu)
+ Au(s)
19
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 + OHE = E0 - 0.059 pH
20
Pourbaix diagram for Mn
21
Example – Group 13
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