1. Chapter 20: Electrochemistry
Chemistry 140 Fall 2002
General Chemistry
Principles and Modern Applications
Petrucci • Harwood • Herring
8th Edition
Chapter 20: Electrochemistry
Philip Dutton
University of Windsor, Canada
N9B 3P4
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2. General Chemistry: Chapter 21
Chemistry 140 Fall 2002
Contents
20-1 Electrode Potentials and Their Measurement
20-2 Standard Electrode Potentials
20-3 Ecell, ΔG, and Keq
20-4 Ecell as a Function of Concentration
20-5 Batteries: Producing Electricity Through Chemical Reactions.
20-6 Corrosion: Unwanted Voltaic Cells
20-7 Electrolysis: Causing Non-spontaneous Reactions to Occur
20-8 Industrial Electolysis Processes
Focus On Membrane Potentials
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General Chemistry: Chapter 21

3. 20-1 Electrode Potentials and Their Measurement
Cu(s) + 2Ag+(aq)
Cu2+(aq) + 2 Ag(s)
Cu(s) + Zn2+(aq)
No reaction
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General Chemistry: Chapter 21

4. An Electrochemical Half Cell
Anode
Cathode
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General Chemistry: Chapter 21

5. An Electrochemical Cell
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General Chemistry: Chapter 21

6. General Chemistry: Chapter 21
Terminology
Electromotive force, Ecell.
The cell voltage or cell potential.
Cell diagram.
Shows the components of the cell in a symbolic way.
Anode (where oxidation occurs) on the left.
Cathode (where reduction occurs) on the right.
Boundary between phases shown by |.
Boundary between half cells (usually a salt bridge) shown by ||.
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General Chemistry: Chapter 21

7. General Chemistry: Chapter 21
Terminology
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) Ecell = V
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General Chemistry: Chapter 21

8. Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) Ecell = 1.103 V
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General Chemistry: Chapter 21

9. General Chemistry: Chapter 21
Terminology
Galvanic cells (Voltaic, or electrochemical).
Produce electricity as a result of spontaneous reactions.
Electrolytic cells.
Non-spontaneous chemical change driven by electricity.
Couple, M|Mn+
A pair of species related by a change in number of e-.
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General Chemistry: Chapter 21

10. 20-2 Standard Electrode Potentials
Cell voltages, the potential differences between electrodes, are among the most precise scientific measurements.
The potential of an individual electrode is difficult to establish.
Arbitrary zero is chosen.
The Standard Hydrogen Electrode (SHE)
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General Chemistry: Chapter 21

11. Standard Hydrogen Electrode
Chemistry 140 Fall 2002
Standard Hydrogen Electrode
2 H+(a = 1) + 2 e- ↔ H2(g, 1 bar) E° = 0 V
Pt|H2(g, 1 bar)|H+(aq, a = 1)
The two vertical lines indicate three phases are present.
For simplicity we usually assume that a = 1 at [H+] = 1 M and replace 1 bar by 1 atm.
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General Chemistry: Chapter 21

12. Standard Electrode Potential, E°
E° defined by international agreement.
The tendency for a reduction process to occur at an electrode.
All ionic species present at a=1 (approximately 1 M).
All gases are at 1 bar (approximately 1 atm).
Where no metallic substance is indicated, the potential is established on an inert metallic electrode (ex. Pt).
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General Chemistry: Chapter 21

13. General Chemistry: Chapter 21
Reduction Couples
Cu2+(1M) + 2 e- → Cu(s) E°Cu/Cu2+ = ?
Pt|H2(g, 1 bar)|H+(a = 1) || Cu2+(1 M)|Cu(s) E°cell = V
anode
cathode
Standard cell potential: the potential difference of a cell formed from two standard electrodes.
E°cell = E°cathode - E°anode
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General Chemistry: Chapter 21

14. Standard Cell Potential
Pt|H2(g, 1 bar)|H+(a = 1) || Cu2+(1 M)|Cu(s) E°cell = V
E°cell = E°cathode - E°anode
E°cell = E°Cu2+/Cu - E°H+/H2
0.340 V = E°Cu2+/Cu - 0 V
E°Cu2+/Cu = V
H2(g, 1 atm) + Cu2+(1 M) → H+(1 M) + Cu(s) E°cell = V
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General Chemistry: Chapter 21

15. Measuring Standard Reduction Potential
anode
cathode
cathode
anode
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General Chemistry: Chapter 21

16. Table 19: Standard reduction potentials
A temperature of K (25 °C). An activity of unity for each pure solid, pure liquid, or for water (solvent).
Legend: (s) – solid; (l) – liquid; (g) – gas; (aq) – aqueous (default for all charged species).  For example Li(s)/Li+ denotes the half reaction Li+ + e- → Li(s), where Li+(aqueous) 
System
E0(V)
Li(s)/Li+
-3.020
K(s)/K+
-2.920
Ba(s)/Ba2+
-2.900
Sr(s)/Sr2+
-2.890
Ca(s)/Ca2+
-2.870
Na(s)/Na+
-2.710
General Chemistry: Chapter 21

17. General Chemistry: Chapter 21
System
E0(V)
H2(g)/2H+
0.000
Sn(s)/Sn4+
0.050
Sn2+/Sn4+
0.150
Cu+/Cu2+
0.153
4OH-/O2(g)+2H2O(l)
0.410
Fe2+/Fe3+
0.760
Ag(s)/Ag+
0.800
2Cl-/Cl2(g)
1.360
Pb2+/Pb4+
1.690
2F-/F2(g)
2.850
General Chemistry: Chapter 21

18. General Chemistry: Chapter 21
20-3 Ecell, ΔG, and Keq
Cells do electrical work.
Moving electric charge.
Faraday constant, F = 96,485 C mol-1
wmaxusef = welec = -nFE
ΔG = -nFE
ΔG° = -nFE°
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General Chemistry: Chapter 21

19. Combining Half Reactions
Fe3+(aq) + 3e- → Fe(s) E°Fe3+/Fe = ?
Fe2+(aq) + 2e- → Fe(s) E°Fe2+/Fe = V
ΔG° = ·F J
Fe3+(aq) + 1e- → Fe2+(aq) E°Fe3+/Fe2+ = V
ΔG° = ·F J
Fe3+(aq) + 3e- → Fe(s)
E°Fe3+/Fe = V
ΔG° = ·F J
-3·F·E° = ΔG° = ·F J
E°Fe3+/Fe = ·F J /(-3·F C) = V = (2 E°Fe2+/Fe + E°Fe3+/Fe2+ )/3 = V
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General Chemistry: Chapter 21

20. General Chemistry: Chapter 21
Spontaneous Change
ΔG < 0 for spontaneous change.
Therefore E°cell > 0 because ΔG°cell = -nFE°cell
E°cell > 0
Reaction proceeds spontaneously as written.
E°cell = 0
Reaction is at equilibrium.
E°cell < 0
Reaction proceeds in the reverse direction spontaneously.
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General Chemistry: Chapter 21

21. The Behavior of Metals Toward Acids
M(s) → M2+(aq) + 2 e- E° = -E°M2+/M
2 H+(aq) + 2 e- → H2(g) E°H+/H2 = 0 V
2 H+(aq) + M(s) → H2(g) + M2+(aq)
E°cell = E°H+/H2 - E°M2+/M = -E°M2+/M
When E°M2+/M < 0, E°cell > 0. Therefore ΔG° < 0.
Metals with negative reduction potentials react with acids
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General Chemistry: Chapter 21

22. Relationship Between E°cell and Keq
ΔG° = -RT ln Keq = -nFE°cell
E°cell = nF RT
ln Keq
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General Chemistry: Chapter 21

23. General Chemistry: Chapter 21
Summary of Thermodynamic, Equilibrium and Electrochemical Relationships.
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General Chemistry: Chapter 21

24. 20-4 Ecell as a Function of Concentration
ΔG = ΔG° +RT ln Q
-nFEcell = -nFEcell° +RT ln Q
Ecell = Ecell° ln Q nF RT
log Q E -4
1.221 -3
1.192 -2
1.162 -1
1.133
1.103 1
1.073 2
1.044 3
1.014 4
0.985
Convert to log10 and calculate constants
Ecell = Ecell° log Q n V
The Nernst Equation:
General Chemistry: Chapter 21

25. General Chemistry: Chapter 21
Example 20-8
Applying the Nernst Equation for Determining Ecell.
What is the value of Ecell for the voltaic cell pictured below and diagrammed as follows?
Pt|Fe2+(0.10 M),Fe3+(0.20 M)||Ag+(1.0 M)|Ag(s)
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General Chemistry: Chapter 21

26. General Chemistry: Chapter 21
Example 20-8
Ecell = Ecell° log Q n V
Ecell = Ecell° log n V
[Fe3+]
[Fe2+] [Ag+]
Ecell = V – V = V
Pt|Fe2+(0.10 M),Fe3+(0.20 M)||Ag+(1.0 M)|Ag(s)
Fe2+(aq) + Ag+(aq) → Fe3+(aq) + Ag (s)
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General Chemistry: Chapter 21

27. General Chemistry: Chapter 21
Chemistry 140 Fall 2002
Concentration Cells
Two half cells with identical electrodes but different ion concentrations.
Pt|H2 (1 atm)|H+(x M)||H+(1.0 M)|H2(1 atm)|Pt(s)
2 H+(1 M) + 2 e- → H2(g, 1 atm)
H2(g, 1 atm) → 2 H+(x M) + 2 e-
2 H+(1 M) → 2 H+(x M)
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General Chemistry: Chapter 21

28. General Chemistry: Chapter 21
Concentration Cells
Ecell = Ecell° log Q n V
2 H+(1 M) → 2 H+(x M)
Ecell = Ecell° log n V x2 12
Ecell = log 2 V x2 1
Ecell = V log x
Ecell = ( V) pH
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General Chemistry: Chapter 21

29. General Chemistry: Chapter 21
Chemistry 140 Fall 2002
Measurement of Ksp
Ag|Ag+(sat’d AgI)||Ag+(0.10 M)|Ag(s)
Ag+(0.100 M) + e- → Ag(s)
Ag(s) → Ag+(sat’d) + e-
Ag+(0.100 M) → Ag+(sat’d M)
Ion concentration difference provides a basis for determining Ksp
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General Chemistry: Chapter 21

30. General Chemistry: Chapter 21
Example 20-10
Using a Voltaic Cell to Determine Ksp of a Slightly Soluble Solute.
With the data given for the reaction on the previous slide, calculate Ksp for AgI.
AgI(s) → Ag+(aq) + I-(aq)
Let [Ag+] in a saturated Ag+ solution be x:
Ag+(0.100 M) → Ag+(sat’d M)
Ecell = Ecell° log Q = n V
Ecell° log
[Ag+]0.10 M soln
[Ag+]sat’d AgI
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General Chemistry: Chapter 21

31. General Chemistry: Chapter 21
Example 20-10
Ecell =
Ecell° log n V
[Ag+]0.10 M soln
[Ag+]sat’d AgI
Ecell =
Ecell° log n V
0.100 x
0.417 =
(log x – log 0.100) 1 V
0.417
log
0.0592
log x =
= -1 – 7.04 = -8.04
x = = 9.1·10-9
Ksp = x2 = 8.3·10-17
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General Chemistry: Chapter 21

32. 20-5 Batteries: Producing Electricity Through Chemical Reactions
Primary Cells (or batteries).
Cell reaction is not reversible.
Secondary Cells.
Cell reaction can be reversed by passing electricity through the cell (charging).
Flow Batteries and Fuel Cells.
Materials pass through the battery which converts chemical energy to electric energy.
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General Chemistry: Chapter 21

33. The Leclanché (Dry) Cell
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General Chemistry: Chapter 21

34. General Chemistry: Chapter 21
Dry Cell
Zn(s) → Zn2+(aq) + 2 e-
Oxidation:
2 MnO2(s) + H2O(l) + 2 e- → Mn2O3(s) + 2 OH-
Reduction:
NH4+ + OH- → NH3(g) + H2O(l)
Acid-base reaction:
NH3 + Zn2+(aq) + Cl- → [Zn(NH3)2]Cl2(s)
Precipitation reaction:
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General Chemistry: Chapter 21

35. General Chemistry: Chapter 21
Alkaline Dry Cell
Reduction:
2 MnO2(s) + H2O(l) + 2 e- → Mn2O3(s) + 2 OH-
Oxidation reaction can be thought of in two steps:
Zn(s) → Zn2+(aq) + 2 e-
Zn2+(aq) + 2 OH- → Zn (OH)2(s)
Zn (s) + 2 OH- → Zn (OH)2(s) + 2 e-
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General Chemistry: Chapter 21

36. Lead-Acid (Storage) Battery
The most common secondary battery
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General Chemistry: Chapter 21

37. General Chemistry: Chapter 21
Lead-Acid Battery
Reduction:
PbO2(s) + 3 H+(aq) + HSO4-(aq)+ 2 e- → PbSO4(s) + 2 H2O(l)
Oxidation:
Pb(s)+ HSO4-(aq)→ PbSO4(s) + H+(aq) + 2 e-
PbO2(s) + Pb(s) + 2 H+(aq) + 2HSO4-(aq) → 2 PbSO4(s) + 2 H2O(l)
E°cell = E°PbO2/PbSO4 - E°PbSO4/Pb = 1.74 V – (-0.28 V) = 2.02 V
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General Chemistry: Chapter 21

38. The Silver-Zinc Cell: A Button Battery
Zn(s),ZnO(s)|KOH(sat’d)|Ag2O(s),Ag(s)
Zn(s) + Ag2O(s) → ZnO(s) + 2 Ag(s) Ecell = 1.8 V
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General Chemistry: Chapter 21

39. The Nickel-Cadmium Cell
Cd(s) + 2 NiO(OH)(s) + 2 H2O(l) → 2 Ni(OH)2(s) + Cd(OH)2(s)
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General Chemistry: Chapter 21

40. General Chemistry: Chapter 21
Fuel Cells
O2(g) + 2 H2O(l) + 4 e- → 4 OH-(aq)
2{H2(g) + 2 OH-(aq) → 2 H2O(l) + 2 e-}
2H2(g) + O2(g) → 2 H2O(l)
E°cell = E°O2/OH- - E°H2O/H2
= V – ( V) = V
 = ΔG°/ ΔH° = 0.83
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General Chemistry: Chapter 21

41. General Chemistry: Chapter 21
Air Batteries
4 Al(s) + 3 O2(g) + 6 H2O(l) + 4 OH- → 4 [Al(OH)4]-(aq)
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General Chemistry: Chapter 21

42. 20-6 Corrosion: Unwanted Voltaic Cells
In neutral solution:
O2(g) + 2 H2O(l) + 4 e- → 4 OH-(aq)
EO2/OH- = V
2 Fe2+(aq) + 4 e- → 2 Fe(s)
EFe/Fe2+ = V
2 Fe(s) + O2(g) + 2 H2O(l) → 2 Fe2+(aq) + 4 OH-(aq)
Ecell = V
In acidic solution:
O2(g) + 4 H+(aq) + 4 e- → 4 H2O (aq) EO2/OH- = V
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General Chemistry: Chapter 21

43. General Chemistry: Chapter 21
Stability of Water pH
2H+/H2
O2(aq)/4OH-
0.000
1.229 1
-0.059
1.170 2
-0.118
1.111 3
-0.178
1.051 4
-0.237
0.992 5
-0.296
0.933 6
-0.355
0.874 7
-0.414
0.815 8
-0.474
0.755 9
-0.533
0.696 10
-0.592
0.637 11
-0.651
0.578 12
-0.710
0.519 13
-0.770
0.459 14
-0.829
0.400
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General Chemistry: Chapter 21

44. General Chemistry: Chapter 21
Corrosion
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General Chemistry: Chapter 21

45. General Chemistry: Chapter 21
Corrosion Protection
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General Chemistry: Chapter 21

46. General Chemistry: Chapter 21
Corrosion Protection
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General Chemistry: Chapter 21

47. 20-7 Electrolysis: Causing Non-spontaneous Reactions to Occur
Galvanic Cell:
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) E = V
Electolytic Cell:
Zn2+(aq) + Cu(s) → Zn(s) + Cu2+(aq) E = V
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General Chemistry: Chapter 21

48. Complications in Electrolytic Cells
Chemistry 140 Fall 2002
Complications in Electrolytic Cells
Overpotential.
Competing reactions.
Non-standard states.
Nature of electrodes.
Overcome interactions a the electrode surface
Hg and H2 overpotential is 1.5 V
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49. Quantitative Aspects of Electrolysis
1 mol e- = C = F
Charge (C) = current (C/s) · time (s)
ne- =
I · t F
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50. 20-8 Industrial Electrolysis Processes
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51. General Chemistry: Chapter 21
Electroplating
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52. General Chemistry: Chapter 21
Chlor-Alkali Process
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General Chemistry: Chapter 21

53. Focus On Membrane Potentials
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General Chemistry: Chapter 21

54. General Chemistry: Chapter 21
Nernst Potential, Δ
For one atom: ∆ = 𝑘 𝑇 𝑧 𝑒 ln⁡( 𝑎 1 𝑎 2 )
Boltzmann constant: k electron charge: e
For one mol
∆ = 𝑅 𝑇 𝑧 𝐹 ln⁡( 𝑎 1 𝑎 2 )
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General Chemistry: Chapter 21

55. General Chemistry: Chapter 21
Ion transports
General Chemistry: Chapter 21

56. Calculate the reversal potential, Δ
Calculate the reversal (Nernst) potential (ΔΦ) for Na+ for a cell in assuming the following ion activities: aint[Na+]=10.18×10−3, aext[Na+]=149.8×10−3, where aint[Na+] is the intracellular activity (≈ concentration in M) of sodium, and  aext[Na+] is the extracellular activity (≈ concentration in M) of sodium.
What is the ΔΦ reversal potential (mV) at 37 ºC?  The Boltzmann constant: k = 1.381×10−23 J/K. The electron's charge: e = 1.602×10−19 C.
General Chemistry: Chapter 21

57. General Chemistry: Chapter 21
Chapter 21 Questions
Develop problem solving skills and base your strategy not on solutions to specific problems but on understanding.
Choose a variety of problems from the text as examples.
Practice good techniques and get coaching from people who have been here before.
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General Chemistry: Chapter 21