1. Chapter 20: Electrochemistry
Chemistry 140 Fall 2002
Chapter 20: Electrochemistry
Juana Mendenhall, Ph.D.
Assistant Professor
Lecture 2 March 15
General Chemistry: Chapter 20
Prentice-Hall © 2007

2. General Chemistry: Chapter 20
Chemistry 140 Fall 2002
Objectives
Describe and illustrate how a voltaic cell operates, with the concepts of electrodes, salt bridges, half-cell equations, net cell reaction and cell diagram.
Describe the standard hydrogen electrode and explain how other standard electrode potentials are related to it.
Using tabulated standard potentials, Eº, to determine E°cell for an oxidation-reduction reaction and predict whether the reaction is spontaneous.
General Chemistry: Chapter 20
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3. An Electrochemical Half Cell
Chemistry 140 Fall 2002
An Electrochemical Half Cell
Anode
Cathode
General Chemistry: Chapter 20
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4. An Electrochemical Cell
Chemistry 140 Fall 2002
An Electrochemical Cell
General Chemistry: Chapter 20
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5. General Chemistry: Chapter 20
Chemistry 140 Fall 2002
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 ||.
General Chemistry: Chapter 20
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6. Terminology Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) Ecell = 1.103 V
Chemistry 140 Fall 2002
Terminology
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
Ecell = V
This typical electrochemical cell was one of the first developed to produce
electrical energy (to run Morse code machines). It is called a Daniell cell.
General Chemistry: Chapter 20
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7. General Chemistry: Chapter 20
Chemistry 140 Fall 2002
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-.
General Chemistry: Chapter 20
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8. General Chemistry: Chapter 20
Chemistry 140 Fall 2002
Example
A galvanic cell consists of a Mg electrode in a 1.0M Mg(NO3)2 solution
and a Ag electrode in a 1.0M AgNO3 solution. Calculate the standard
emf of this electrochemical cell at 25º C.
Answer
Mg2+(aq) + 2e Mg(s) Eº = V
Ag+(aq) + e Ag(s) °E = V
Ag+ will oxidize the Mg2+ (diagonal rule):
Anode: Mg(s) Mg2+(aq) + 2e-
Cathode: 2Ag+(aq) + 2e Ag(s)
Overall: Mg (s) + 2Ag+(aq) Mg2+(aq) Ag(s)
General Chemistry: Chapter 20
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9. General Chemistry: Chapter 20
Chemistry 140 Fall 2002
Example cont.
Eº = Eºmg/mg2+ + EºAg+/Ag
= 2.37 V V
= 3.17 V
The sign of the emf of the cell allows us to predict the spontaneity of a
redox rxn. Recall that under steady-state conditions for rxns & products,
the redox rxn is spontaneous in the forward direction if the standard
emf of the cell is positive. If it is negative, Eºcell does not mean the rxn
will not occur, it means the the equilibrium will shift to the left.
General Chemistry: Chapter 20
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10. Standard Electrode Potentials
Chemistry 140 Fall 2002
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)
General Chemistry: Chapter 20
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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+(a = 1)
General Chemistry: Chapter 20
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12. Standard Electrode Potential, E°
Chemistry 140 Fall 2002
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).
General Chemistry: Chapter 20
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13. General Chemistry: Chapter 20
Chemistry 140 Fall 2002
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
General Chemistry: Chapter 20
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14. Standard Cell Potential
Chemistry 140 Fall 2002
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
General Chemistry: Chapter 20
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15. General Chemistry: Chapter 20
Chemistry 140 Fall 2002
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°
General Chemistry: Chapter 20
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16. General Chemistry: Chapter 20
Chemistry 140 Fall 2002
Example
Calculate the standard free-energy for the following 25ºC:
2Au(s) + 3Ca2+ (aq) Au3+(aq) Ca(s)
Answer: First we break up the overall rxn into half-rxns:
Oxidation: Au(s) Au3+(aq) + 6e-
Reduction: 3Ca2+ (aq) + 6e Ca(s)
Eº = EºAu/Au3+ + EºCa2+/Ca
= V V
= V
Gº = -nFEº
= - (6 mol)(96,500 J/V*mol)(-4.37 V)
= x 103 kJ
General Chemistry: Chapter 20
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17. General Chemistry: Chapter 20
Chemistry 140 Fall 2002
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.
General Chemistry: Chapter 20
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18. Combining Half Reactions
Chemistry 140 Fall 2002
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
General Chemistry: Chapter 20
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