1. Chapter 20: Electrochemistry
Chemistry 140 Fall 2002
CHEMISTRY
Ninth
Edition
GENERAL
Principles and Modern Applications
Petrucci • Harwood • Herring • Madura
Chapter 20: Electrochemistry
General Chemistry: Chapter 20
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2. General Chemistry: Chapter 20
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 Electrolysis Processes
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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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4. An Electrochemical Half Cell
Anode
Cathode
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5. An Electrochemical Cell
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6. General Chemistry: Chapter 20
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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7. Terminology Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) Ecell = 1.103 V
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8. General Chemistry: Chapter 20
Terminology
Galvanic cells.
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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9. 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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10. 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+(a = 1)
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11. 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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12. General Chemistry: Chapter 20
Reduction Couples
Cu2+(1M) + 2 e- → Cu(s) E°Cu2+/Cu = ?
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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13. 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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14. Measuring Standard Reduction Potential
anode
cathode
cathode
anode
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15. Standard Reduction Potentials
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16. General Chemistry: Chapter 20
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17. General Chemistry: Chapter 20
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18. General Chemistry: Chapter 20
20-3 Ecell, ΔG, and Keq
Cells do electrical work.
Moving electric charge.
Faraday constant, F = 96,485 C mol-1
elec = -nFE
Michael Faraday
ΔG = -nFE
ΔG° = -nFE°
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19. Combining Half Reactions
Fe3+(aq) + 3e- → Fe(s) E°Fe3+/Fe = ?
Fe2+(aq) + 2e- → Fe(s) E°Fe2+/Fe = V
ΔG° = J
Fe3+(aq) + 3e- → Fe2+(aq) E°Fe3+/Fe2+ = V
ΔG° = J
Fe3+(aq) + 3e- → Fe(s)
E°Fe3+/Fe = V
ΔG° = V
ΔG° = V = -nFE°
E°Fe3+/Fe = V /(-3F) = V
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20. General Chemistry: Chapter 20
Spontaneous Change
ΔG < 0 for spontaneous change.
Therefore E°cell > 0 because ΔGcell = -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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21. The Behavior or 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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22. Relationship Between E°cell and Keq
ΔG° = -RT ln Keq = -nFE°cell
E°cell = nF RT
ln Keq
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23. General Chemistry: Chapter 20
Summary of Thermodynamic, Equilibrium and Electrochemical Relationships.
General Chemistry: Chapter 20
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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
Convert to log10 and calculate constants.
Ecell = Ecell° log Q n V
The Nernst Equation:
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25. General Chemistry: Chapter 20
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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26. General Chemistry: Chapter 20
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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27. General Chemistry: Chapter 20
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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28. General Chemistry: Chapter 20
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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29. General Chemistry: Chapter 20
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
General Chemistry: Chapter 20
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30. General Chemistry: Chapter 20
EXAMPLE 20-10
Using a Voltaic Cell to Determine Ksp of a Slightly Soluble Solute. With the date 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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31. General Chemistry: Chapter 20
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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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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33. The Leclanché (Dry) Cell
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34. General Chemistry: Chapter 20
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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35. General Chemistry: Chapter 20
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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36. Lead-Acid (Storage) Battery
The most common secondary battery.
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37. General Chemistry: Chapter 20
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) + HSO4-(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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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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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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40. General Chemistry: Chapter 20
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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41. General Chemistry: Chapter 20
Air Batteries
4 Al(s) + 3 O2(g) + 6 H2O(l) + 4 OH- → 4 [Al(OH)4](aq)
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42. 20-6 Corrosion: Unwanted Voltaic Cells
In neutral solution:
O2(g) + 2 H2O(l) + 4 e- → 4 OH-(aq)
EO2/OH- = V
2 Fe(s) → 2 Fe2+(aq) + 4 e-
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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43. General Chemistry: Chapter 20
Corrosion
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44. General Chemistry: Chapter 20
Corrosion Protection
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45. General Chemistry: Chapter 20
Corrosion Protection
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46. 20-7 Electrolysis: Causing Non-spontaneous Reactions to Occur
Galvanic Cell:
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) EO2/OH- = V
Electolytic Cell:
Zn2+(aq) + Cu(s) → Zn(s) + Cu2+(aq) EO2/OH- = V
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47. Predicting Electrolysis Reaction
An Electrolytic Cell
e- is the reverse of the voltaic cell.
Battery must have a voltage in excess of V in order to force the non-spontaneous reaction.
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48. Complications in Electrolytic Cells
Chemistry 140 Fall 2002
Complications in Electrolytic Cells
Overpotential.
Competing reactions.
Non-standard states.
Nature of electrodes.
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49. Quantitative Aspects of Electrolysis
1 mol e- = C
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 20
Chlor-Alkali Process
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