1. Chapter 17: Electrochemistry
12/26/2017
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

2. Chapter 17: Electrochemistry
12/26/2017
Galvanic Cells
Electrochemistry: The area of chemistry concerned with the interconversion of chemical and electrical energy.
Galvanic (Voltaic) Cell: A spontaneous chemical reaction which generates an electric current.
Electrolytic Cell: An electric current which drives a nonspontaneous reaction.
Galvanic and electrolytic cells are the reverse of each other.
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

3. Chapter 17: Electrochemistry
12/26/2017
Galvanic Cells
Zn2+(aq) + Cu(s)
Zn(s) + Cu2+(aq)
Oxidation half-reaction:
Zn2+(aq) + 2e-
Zn(s)
Reduction half-reaction:
Cu(s)
Cu2+(aq) + 2e-
Balancing redox reactions was covered in Ch 4 (Reactions in Aqueous Solution). It’s common for people to postpone redox balancing until now.
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

4. Chapter 17: Electrochemistry
12/26/2017
Galvanic Cells
Zn2+(aq) + Cu(s)
Zn(s) + Cu2+(aq)
Be careful about assumptions. Don’t always place the same electrode on the same side. For instance, in the past, textbooks would commonly placed the anode on the left side which led some students to believe that the anode always belongs on the left.
Students often are confused as the purpose and need of a salt bridge.
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5. Chapter 17: Electrochemistry
12/26/2017
Galvanic Cells
Anode:
The electrode where oxidation occurs.
The electrode where electrons are produced.
Is what anions migrate toward.
Has a negative sign.
The wire convention is used for the sign at the anode and cathode.
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Copyright © 2008 Pearson Prentice Hall, Inc.

6. Chapter 17: Electrochemistry
12/26/2017
Galvanic Cells
Anode:
The electrode where oxidation occurs.
The electrode where electrons are produced.
Is what anions migrate toward.
Has a negative sign.
Cathode:
The electrode where reduction occurs.
The electrode where electrons are consumed.
Is what cations migrate toward.
Has a positive sign.
“Anode” and “oxidation” both begin with vowels.
“Cathode” and “reduction” both begin with consonants.
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

7. Chapter 17: Electrochemistry
12/26/2017
Galvanic Cells
Anode half-reaction:
Zn2+(aq) + 2e-
Zn(s)
Cathode half-reaction:
Cu(s)
Cu2+(aq) + 2e-
Overall cell reaction:
Zn2+(aq) + Cu(s)
Zn(s) + Cu2+(aq)
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

8. Shorthand Notation for Galvanic Cells
Chapter 17: Electrochemistry
12/26/2017
Shorthand Notation for Galvanic Cells
Anode half-reaction:
Zn2+(aq) + 2e-
Zn(s)
Cathode half-reaction:
Cu(s)
Cu2+(aq) + 2e-
Overall cell reaction:
Zn2+(aq) + Cu(s)
Zn(s) + Cu2+(aq)
Salt bridge
Anode half-cell
Cathode half-cell
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
Electron flow
Phase boundary
Phase boundary
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Copyright © 2008 Pearson Prentice Hall, Inc.

9. Cell Potentials and Free-Energy Changes for Cell Reactions
Chapter 17: Electrochemistry
12/26/2017
Cell Potentials and Free-Energy Changes for Cell Reactions
Electromotive Force (emf): The force or electrical potential that pushes the negatively charged electrons away from the anode (- electrode) and pulls them toward the cathode (+ electrode).
It is also called the cell potential (E) or the cell voltage.
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

10. Cell Potentials and Free-Energy Changes for Cell Reactions
Chapter 17: Electrochemistry
12/26/2017
Cell Potentials and Free-Energy Changes for Cell Reactions
1 J = 1 C x 1 V
joule
SI unit of energy
volt
SI unit of electric potential
coulomb
Electric charge
1 coulomb is the amount of charge transferred when a current of 1 ampere flows for 1 second.
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

11. Cell Potentials and Free-Energy Changes for Cell Reactions
Chapter 17: Electrochemistry
12/26/2017
Cell Potentials and Free-Energy Changes for Cell Reactions
faraday or Faraday constant
the electric charge on 1 mol of electrons
96,5000 C/mol e-
DG = -nFE or
DG° = -nFE°
free-energy change
The “°” refers to standard free-energy change and standard cell potential.
cell potential
number of moles of electrons transferred in the reaction
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

12. Cell Potentials and Free-Energy Changes for Cell Reactions
Chapter 17: Electrochemistry
12/26/2017
Cell Potentials and Free-Energy Changes for Cell Reactions
The standard cell potential at 25 °C is 0.10 V for the reaction:
Zn2+(aq) + Cu(s)
Zn(s) + Cu2+(aq)
Calculate the standard free-energy change for this reaction at 25 °C.
DG° = -nFE°
The sign for a spontaneous reaction for free-energy: -
The sign for a spontaneous reaction for cell potential: +
mol e-
96,500 C
1 C V
1 J
1000 J
1 kJ
= -(2 mol e-)
(1.10 V)
DG° = -212 kJ
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Copyright © 2008 Pearson Prentice Hall, Inc.

13. Standard Reduction Potentials
Chapter 17: Electrochemistry
12/26/2017
Standard Reduction Potentials
2H1+(aq) + Cu(s)
H2(g) + Cu2+(aq)
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14. Standard Reduction Potentials
Chapter 17: Electrochemistry
12/26/2017
Standard Reduction Potentials
Anode half-reaction:
2H1+(aq) + 2e-
H2(g)
Cathode half-reaction:
Cu(s)
Cu2+(aq) + 2e-
Overall cell reaction:
2H1+(aq) + Cu(s)
H2(g) + Cu2+(aq)
The standard potential of a cell is the sum of the standard half-cell potentials for oxidation at the anode and reduction at the cathode:
Individual half-cell potentials can’t be measured. They must be measured in pairs.
E°cell = E°ox + E°red
The measured potential for this cell: E°cell = 0.34 V
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Copyright © 2008 Pearson Prentice Hall, Inc.

15. Standard Reduction Potentials
Chapter 17: Electrochemistry
12/26/2017
Standard Reduction Potentials
The standard hydrogen electrode (S.H.E.) has been chosen to be the reference electrode.
Potentials can only be measured in pairs which is why we need to establish a reference point for our half-cell potentials. Much like what was done for standard enthalpies and standard free energies.
Copyright © 2008 Pearson Prentice Hall, Inc.

16. Standard Reduction Potentials
Chapter 17: Electrochemistry
12/26/2017
Standard Reduction Potentials
The standard hydrogen electrode (S.H.E.) has been chosen to be the reference electrode.
H2(g, 1 atm)
2H1+(aq, 1 M) + 2e-
E°ox = 0 V
E°red = 0 V
2H1+(aq, 1 M) + 2e-
H2(g, 1 atm)
The half-cell potentials for the S.H.E. have been defined to be 0 V. Much like the atomic mass of carbon-12 has been defined to 12 u.
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Copyright © 2008 Pearson Prentice Hall, Inc.

17. Standard Reduction Potentials
Chapter 17: Electrochemistry
12/26/2017
Standard Reduction Potentials
Anode half-reaction:
2H1+(aq) + 2e-
H2(g)
Cathode half-reaction:
Cu(s)
Cu2+(aq) + 2e-
Overall cell reaction:
2H1+(aq) + Cu(s)
H2(g) + Cu2+(aq)
E°cell = E°ox + E°red
0.34 V = 0 V + E°red
Recall that the 0.34 V was measured for the cell with the half-cells as described above. The oxidation half-cell potential is defined to be 0 V since it is the S.H.E.
Note also that the “ox” and “red” subscripts for oxidation and reduction have been dropped. While this can be very confusing to students, it is common usage.
A standard reduction potential can be defined:
Cu(s)
Cu2+(aq) + 2e-
E° = 0.34 V
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

18. Standard Reduction Potentials
Chapter 17: Electrochemistry
12/26/2017
Standard Reduction Potentials
H2(g) + Zn2+(aq)
2H1+(aq) + Zn(s)
Copyright © 2008 Pearson Prentice Hall, Inc.

19. Standard Reduction Potentials
Chapter 17: Electrochemistry
12/26/2017
Standard Reduction Potentials
Anode half-reaction:
H2(g)
2H1+(aq) + 2e-
Cathode half-reaction:
Zn2+(aq) + 2e-
Zn(s)
Overall cell reaction:
H2(g) + Zn2+(aq)
2H1+(aq) + Zn(s)
E°cell = E°ox + E°red
0.76 V = E°ox + 0 V
The half-cell potential sign flips for an oxidation potential (the table has reduction potentials).
Zn2+(aq) + 2e-
Zn(s)
E° = 0.76 V
As a standard reduction potential:
Zn(s)
Zn2+(aq) + 2e-
E° = V
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

20. Standard Reduction Potentials
Chapter 17: Electrochemistry
12/26/2017
Standard Reduction Potentials
Copyright © 2008 Pearson Prentice Hall, Inc.

21. Using Standard Reduction Potentials
Chapter 17: Electrochemistry
12/26/2017
Using Standard Reduction Potentials
Zn2+(aq) + 2e-
Zn(s)
E° = -(-0.76 V)
Cu(s)
Cu2+(aq) + 2e-
E° = 0.34 V
E° = 1.10 V
Zn2+(g) + Cu(s)
Zn(s) + Cu2+(aq)
Ag(s)]
2 x [Ag1+(aq) + e-
E° = 0.80 V
Cu2+(aq) + 2e-
Cu(s)
E° = V
E° = 0.46 V
2Ag2+(g) + Cu(s)
2Ag(s) + Cu2+(aq)
Half-cell potentials are intensive properties.
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Copyright © 2008 Pearson Prentice Hall, Inc.

22. Chapter 17: Electrochemistry
12/26/2017
The Nernst Equation
DG = DG° + RT ln Q
Using: DG = -nFE and DG° = -nFE°
ln Q nF RT
E = E° -
Nernst Equation: or
log Q nF
2.303RT
E = E° -
n is the number of moles of transferred electrons from the balanced reaction (balanced number of electrons from the half-reactions). Also, just the number is used (“2” instead of “2 mol e-”). or
log Q n V
E = E° -
in volts, at 25°C
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Copyright © 2008 Pearson Prentice Hall, Inc.

23. Chapter 17: Electrochemistry
12/26/2017
The Nernst Equation
Consider a galvanic cell that uses the reaction:
Cu2+(aq) + 2Fe2+(aq)
Cu(s) + 2Fe3+(aq)
What is the potential of a cell at 25 °C that has the following ion concentrations?
[Fe3+] = 1.0 x 10-4 M
[Cu2+] = 0.25 M
[Fe2+] = 0.20 M
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

24. Chapter 17: Electrochemistry
12/26/2017
The Nernst Equation
log Q n V
E = E° -
Calculate E°:
Cu2+(aq) + 2e-
Cu(s)
E° = V
Fe2+(aq)
Fe3+(aq) + e-
E° = 0.77 V
E°cell = V V = 0.43 V
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

25. Chapter 17: Electrochemistry
12/26/2017
The Nernst Equation
log Q n V
E = E° -
Calculate E: n V
[Fe3+]2
[Cu2+][Fe2+]2
E = E° -
log 2 V
(1.0 x 10-4)2
(0.25)(0.20)2
= 0.43 V -
log
E = 0.25 V
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

26. Standard Cell Potentials and Equilibrium Constants
Chapter 17: Electrochemistry
12/26/2017
Standard Cell Potentials and Equilibrium Constants
Using
DG° = -nFE°
and
DG° = -RT ln K
-nFE° = -RT ln K nF RT nF
2.303 RT
E° =
ln K =
log K
DG° = -RT ln K came from Ch 16 (Thermodynamics: Entropy, Free Energy, and Equilibrium).
The mathematical reduction from 2.303RT/nF to V/n is the same as what was done for the Nernst equation. n V
E° =
log K
in volts, at 25°C
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

27. Standard Cell Potentials and Equilibrium Constants
Chapter 17: Electrochemistry
12/26/2017
Standard Cell Potentials and Equilibrium Constants
Copyright © 2008 Pearson Prentice Hall, Inc.

28. Standard Cell Potentials and Equilibrium Constants
Chapter 17: Electrochemistry
12/26/2017
Standard Cell Potentials and Equilibrium Constants
Three methods to determine equilibrium constants:
K =
[A]a[B]b
[C]c[D]d
K from concentration data: RT
-DG°
ln K =
K from thermochemical data:
ln K nF RT
E° =
K from electrochemical data: or RT
nFE°
ln K =
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

29. Chapter 17: Electrochemistry
12/26/2017
Batteries
Lead Storage Battery
Anode:
PbSO4(s) + H1+(aq) + 2e-
Pb(s) + HSO41-(aq)
Cathode:
PbSO4(s) + 2H2O(l)
PbO2(s) + 3H1+(aq) + HSO41-(aq) + 2e-
Overall:
2PbSO4(s) + 2H2O(l)
Pb(s) + PbO2(s) + 2H1+(aq) + 2HSO41-(aq)
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30. Chapter 17: Electrochemistry
12/26/2017
Batteries
Dry-Cell Batteries
Leclanché cell
Anode:
Zn2+(aq) + 2e-
Zn(s)
Cathode:
Mn2O3(s) + 2NH3(aq)+ H2O(l)
2MnO2(s) + 2NH41+(aq) + 2e-
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Copyright © 2008 Pearson Prentice Hall, Inc.

31. Chapter 17: Electrochemistry
12/26/2017
Batteries
Dry-Cell Batteries
Alkaline cell
Anode:
ZnO(s) + H2O(l) + 2e-
Zn(s) + 2OH1-(aq)
Cathode:
Mn2O3(s) + 2OH1-(aq)
2MnO2(s) + H2O(l) + 2e-
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

32. Chapter 17: Electrochemistry
12/26/2017
Batteries
Nickel-Cadmium (“ni-cad”) Batteries
Anode:
Cd(OH)2(s) + 2e-
Cd(s) + 2OH1-(aq)
Cathode:
Ni(OH)2(s) + OH1-(aq)
NiO(OH)(s) + H2O(l) + e-
Nickel-Metal Hydride (“NiMH”) Batteries
Anode:
M(s) + H2O(l) + e-
MHab(s) + OH1-(aq)
Cathode:
Ni(OH)2(s) + OH1-(aq)
NiO(OH)(s) + H2O(l) + e-
Overall:
M(s) + Ni(OH)2(s)
MHab(s) + NiO(OH)(s)
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Copyright © 2008 Pearson Prentice Hall, Inc.

33. Chapter 17: Electrochemistry
12/26/2017
Batteries
Lithium and Lithium Ion Batteries
Lithium
Anode:
xLi1+(soln) + xe-
xLi(s)
Cathode:
LixMnO2(s)
MnO2(s) + xLi1+(soln) + xe-
Lithium Ion
Notice where lithium is on the standard reduction table.
Anode:
xLi1+(soln) + 6C(s) + xe-
LixC6(s)
Cathode:
LiCoO2(s)
Li1-xCoO2(s) + xLi1+(soln) + xe-
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Copyright © 2008 Pearson Prentice Hall, Inc.

34. Chapter 17: Electrochemistry
12/26/2017
Fuel Cells
Hydrogen-Oxygen Fuel Cell
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Copyright © 2008 Pearson Prentice Hall, Inc.

35. Chapter 17: Electrochemistry
12/26/2017
Corrosion
Corrosion: The oxidative deterioration of a metal.
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Copyright © 2008 Pearson Prentice Hall, Inc.

36. Chapter 17: Electrochemistry
12/26/2017
Corrosion
Prevention of Corrosion
For some metals, oxidation protects the metal (aluminum, chromium, magnesium, titanium, zinc, and others). For other metals, there are two main techniques.
Iron rust doesn’t protect iron since the rust is too porous.
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Copyright © 2008 Pearson Prentice Hall, Inc.

37. Chapter 17: Electrochemistry
12/26/2017
Corrosion
Prevention of Corrosion
Galvanization: The coating of iron with zinc.
Even if the zinc coating is scratched, the iron is still protected since the zinc will be selectively oxidized.
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Copyright © 2008 Pearson Prentice Hall, Inc.

38. Chapter 17: Electrochemistry
12/26/2017
Corrosion
Prevention of Corrosion
Galvanization: The coating of iron with zinc.
When some of the iron is oxidized (rust), the process is reversed since zinc will reduce Fe2+ to Fe:
Fe(s)
Fe2+(aq) + 2e-
E° = V
Fe2+ is higher up in the reduction table so it will be selectively reduced when in contact with zinc.
Zn(s)
Zn2+(aq) + 2e-
E° = V
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Copyright © 2008 Pearson Prentice Hall, Inc.

39. Chapter 17: Electrochemistry
12/26/2017
Corrosion
Prevention of Corrosion
Cathodic Protection: Instead of coating the entire surface of the first metal with a second metal, the second metal is placed in electrical contact with the first metal:
Anode:
Mg2+(aq) + 2e-
Mg(s)
E° = 2.37 V
Cathode:
2H2O(l)
O2(g) + 4H1+(aq) + 4e-
Magnesium is farther down in the reduction table (on the right) so it is more likely to be oxidized than zinc.
Zinc is used as a sacrificial anode for iron sea walls for hurricane protection (an iron wall placed into salt water for a long period of time is not a good idea from a corrosion perspective!). Blocks of solid zinc (sacrificial anodes) are attached to the iron and spaced at regular intervals. The sacrificial anodes are monitored over time and replaced when needed.
E° = 1.23 V
Attaching a magnesium stake to iron will corrode the magnesium instead of the iron. Magnesium acts as a sacrificial anode.
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

40. Electrolysis and Electrolytic Cells
Chapter 17: Electrochemistry
12/26/2017
Electrolysis and Electrolytic Cells
Electrolysis: The process of using an electric current to bring about chemical change.
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

41. Electrolysis and Electrolytic Cells
Chapter 17: Electrochemistry
12/26/2017
Electrolysis and Electrolytic Cells
Electrolysis of Molten Sodium Chloride
Anode:
Cl2(g) + 2e-
2Cl1-(l)
Cathode:
2Na(l)
2Na1+(l) + 2e-
Overall:
2Na(l) + Cl2(g)
2Na1+(l) + 2Cl1-(l)
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42. Electrolysis and Electrolytic Cells
Chapter 17: Electrochemistry
12/26/2017
Electrolysis and Electrolytic Cells
Electrolysis of Aqueous Sodium Chloride
Anode:
Cl2(g) + 2e-
2Cl1-(aq)
Cathode:
H2(g) + 2OH1-(aq)
2H2O(l) + 2e-
Overall:
Cl2(g) + H2(g) + 2OH1-(aq)
2Cl1-(l) + 2H2O(l)
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

43. Electrolysis and Electrolytic Cells
Chapter 17: Electrochemistry
12/26/2017
Electrolysis and Electrolytic Cells
Electrolysis of Water
Anode:
O2(g) + 4H1+(aq) + 4e-
2H2O(l)
Cathode:
2H2(g) + 4OH1-(aq)
4H2O(l) + 4e-
Overall:
2H2(g) + O2(g) + 4H1+ + 4OH1-(aq)
6H2O(l)
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

44. Commercial Applications of Electrolysis
Chapter 17: Electrochemistry
12/26/2017
Commercial Applications of Electrolysis
Down’s Cell for the Production of Sodium Metal
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Copyright © 2008 Pearson Prentice Hall, Inc.

45. Commercial Applications of Electrolysis
Chapter 17: Electrochemistry
12/26/2017
Commercial Applications of Electrolysis
A Membrane Cell for Electrolytic Production of Cl2 and NaOH
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Copyright © 2008 Pearson Prentice Hall, Inc.

46. Commercial Applications of Electrolysis
Chapter 17: Electrochemistry
12/26/2017
Commercial Applications of Electrolysis
Hall-Heroult Process for the Production of Aluminum
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

47. Commercial Applications of Electrolysis
Chapter 17: Electrochemistry
12/26/2017
Commercial Applications of Electrolysis
Electrorefining of copper metal
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Copyright © 2008 Pearson Prentice Hall, Inc.

48. Quantitative Aspects of Electrolysis
Chapter 17: Electrochemistry
12/26/2017
Quantitative Aspects of Electrolysis
Charge(C) = Current(A) x Time(s)
96,500 C
1 mol e-
Moles of e- = Charge(A) x
Faraday constant
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.