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CH5. Oxidation and Reduction

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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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Iron smelting

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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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Standard cell and potentials

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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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Frost diagrams

prop to -G

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Frost diagrams

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Frost diagram for N

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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

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Example – Group 13