1. Chapter 18 Electrochemistry

2. Oxidation and Reduction
Oxidation is the loss of electrons by a chemical process.
When sodium forms a compound, Na+ is formed. Sodium is oxidized.
Reduction is the gain of electrons by a chemical process.
When Cl- ions are formed from elemental chlorine, chlorine is reduced. 1

3. Oxidation-Reduction (“Redox”)
An oxidation-reduction reaction, or redox reaction, is one in which electrons are transferred from one species to another.
In every redox reaction, at least one species is oxidized and at least one species is reduced.
2Na(s) + Cl2(g) → 2NaCl(s) is a redox reaction because Na is oxidized and Cl is reduced. 1

4. Oxidizing and Reducing Agents
An oxidizing agent is the reactant that accepts electrons, causing an oxidation to occur.
The oxidizing agent is reduced.
A reducing agent is the reactant that supplies electrons, causing a reduction to occur.
The reducing agent is oxidized.
In the reaction of sodium with chlorine, Na is the reducing agent and Cl2 is the oxidizing agent. 1

5. Half-reactions
In a half-reaction, either the oxidation or reduction part of a redox reaction is given, showing the electrons explicitly.
Half-reactions emphasize the transfer of electrons in a redox reaction.
For 2Na(s) + Cl2(g) → 2NaCl(s) :
Na → Na+ + 1e- oxidation half-reaction
Cl2 + 2e- → 2Cl- reduction half-reaction 1

6. Oxidation States
The oxidation state is the charge on the monatomic ion, or the charge on an atom when the shared electrons are assigned to the more electronegative atom.
Electron pairs shared by atoms of the same element are divided equally.
In CaCl2, an ionic compound:
calcium has an oxidation state of +2.
chlorine has an oxidation state of -1. 1

7. Review
Assign oxidation numbers to each atom in the following substances.
(a) PF3 (b) CO (c) NH4Cl 1

8. Review Balance the following equation in acid solution:
Cr2O72- + C2H5OH → Cr3+ + CO2 7

9. Review Balance the following equation in basic solution:
Zn + ClO- → Zn(OH)42- + Cl- 7

10. Voltaic Cells
A voltaic cell (also known as a galvanic cell) is an apparatus that produces electrical energy directly from a redox reaction.
All voltaic cells depend on redox reactions. 7

11. Voltaic Cell - Diagram 1

12. Voltaic Cells
The oxidation half-cell contains the reaction Zn(s) → Zn2+(aq) + 2e-.
The reduction half-cell contains the reaction Cu2+(aq) + 2e- → Cu(s).
Since the half-cells are physically separated, electrons must travel from one side to another through a connecting wire. 1

13. Voltaic Cell Conventions
Assume that oxidation occurs in the left half-cell and reduction occurs in the right half-cell.
Zn(s) → Zn2+(aq) + 2e- occurs in left.
Cu2+(aq) + 2e- → Cu(s) occurs in right.
A voltmeter is connected to the wire coming out of each metal electrode.
If the measured voltage is positive, the chemical reaction is spontaneous as written. 1

14. Electrodes and Half-Cells
Redox reactions not involving metals can be used in half-cells by using an inert electrode, like gold, platinum, or carbon to provide electrical contact. Examples of half-reactions include:
two soluble ions
Fe3+(aq) + e- → Fe2+(aq)
gas-ion reactions
Cl2(g) + e- → 2Cl-(aq)
insoluble salts
AgCl(s) + e- → Ag(s) + Cl-(aq) 1

15. Example of Inert-Electrode Half-Cell1

16. The Hydrogen Electrode1

17. Electrical Potential
Electromotive force or emf (E) is the electrical driving force that “pushes” electrons from the oxidation half-cell to the reduction half-cell.
Cell potential is the potential energy per unit charge that is characteristic of each half-cell reaction. It is measured in volts (V).
1 volt = 1 joule/coulomb (1 V = 1 J/C) 1

18. Standard Potentials
The overall potential of a cell depends on the concentrations of the species in the reaction.
Standard potential, Ecell, is the cell potential when each species in the reaction is present in its standard state.
Solids, liquids, and gases are in their pure state at 1 atm pressure.
Solutes have 1 M concentration. 1

19. Potentials of Voltaic Cells
For the reaction
Zn(s) + Cu2+(aq) → Cu(s) + Zn2+(aq)
the Ecell is V.
A positive cell potential means that the reaction is spontaneous in the forward direction. 1

20. Additivity of Cell Potentials
Cell potentials are additive:
Cu(s) + 2Ag+(1 M) → 2Ag(s) + Cu2+(1 M) V
Zn(s) + Cu2+(1 M) → Zn2+(1 M) + Cu(s) V
Zn(s) + 2Ag+(1 M) → 2Ag(s) + Zn2+(1 M)
E°cell =+1.10 V V = V 1

21. Standard Reduction Potential
The standard reduction potential of a half-reaction is the potential of the reduction reaction relative to the standard hydrogen electrode as the oxidation. 1

22. Standard Reduction Potentials Table
Half-reactions written as reductions, where electrons are reactants.
Standard reduction potential for
2H+(aq,1 M) + 2e- → H2(g, 1 atm)
is set to V.
If a reduction reaction is reversed to make oxidation, change sign on reduction potential.
Ag+(aq, 1 M) + e- → Ag(s) E = 0.80 V
Ag(s) → Ag+(aq, 1 M) + e- E =-0.80 V 1

23. Standard Reduction Potentials Table
Reduction Half-Reaction E (V)
F2(g) + 2e- → 2F-(aq)
Ag+(aq) + e- → Ag(s)
Fe3+(aq) + e- → Fe2+(aq) 0.77
Sn4+(aq) + 2e- → Sn2+(aq) 0.15
2H+(aq) + 2e- → H2(g)
Co2+(aq) + 2e- → Co(s)
Fe2+(aq) + 2e- → Fe(s)
Zn2+(aq) + 2e- → Zn(s)
Mg2+(aq) + 2e- → Mg(s) 1

24. Calculating Cell Potential
The standard potential of a cell reaction is given by
Ecell = Ered + Eox
Both half-reactions must transfer the same number of electrons.
Multiplying the coefficients of a half-reaction to balance the electrons does NOT change the potential.
Fe3+(aq) + e- → Fe2+(aq) E = 0.77 V
2Fe3+(aq) + 2e- → 2Fe2+(aq) E = 0.77 V 1

25. Example: Using Standard Potentials
Calculate the potential of the cell reaction
2Fe3+ + 2I- → 2Fe2+ + I2
from the potentials in Table 18.1.
Answer: E = 0.23 V 1

26. Test Your Skill
Write the spontaneous cell reaction and calculate Ecell for the voltaic cell made up from the two half-reactions below.
Sn4+(aq) + 2e- → Sn2+(aq) E = 0.15 V
Mg2+(aq) + 2e- → Mg(s) E = V 1

27. Test Your Skill
Write the spontaneous cell reaction and calculate Ecell for the voltaic cell made up from the two half-reactions below.
Sn4+(aq) + 2e- → Sn2+(aq) E = 0.15 V
Mg2+(aq) + 2e- → Mg(s) E = V
Answer:
Mg(s) + Sn4+(aq) → Mg2+(aq) + Sn2+(aq)
Ecell = 2.52 V 1

28. Activity Series
An activity series arranges half-reactions in order of decreasing potential. 1

29. Use of Activity Series
Species with large positive reduction potentials are oxidizing agents – they oxidize species below them in the activity series.
Species with large negative reduction potentials are reducing agents – they reduce species above them in the activity series. 1

30. Activity Series
What happens when Fe is added to 1 M solutions of Zn(NO3)2 and Co(NO3)2?
Co2+(aq) + 2e- → Co(s) E = V
Fe2+(aq) + 2e- → Fe(s) E = V
Zn2+(aq) + 2e- → Zn(s) E = V
Fe reacts with Co2+ but not with Zn2+. 1

31. Test Your Skill
From the data in Table 18.1, select a metal that reduces Ag+(aq) but does not reduce AgCl(s). Write the reaction and calculate its standard potential. 1

32. Cell Potentials and DG DG = -nFE where DG = free energy change;
n = number of electrons transferred;
F = Faraday constant, 96,485 C/mol e-;
E = cell potential. 1

33. Cell Potentials and DG Under standard conditions, DG = -nFE.
This means that spontaneous reactions have positive cell potentials. 1

34. Calculating DG Calculate DG for the reaction below.
AgCl(s) → Ag+(aq) + Cl-(aq)
using the following information:
AgCl(s) + e- ® Ag(s) + Cl-(aq) E° = V
Ag(s) ® Ag+(aq) + e- E° = V 1

35. Test Your Skill
Use the data in Appendix H to calculate DG for the following reaction.
2Na(s) + 2H2O(l) → 2Na+(aq) + 2OH-(aq) + H2(g) 1

36. Relation of E to Keq DG° = -RT ln Keq DG° = -nFE° -nFE° = -RT ln Keq
E° = ln Keq = log Keq
At 298 K, E° = log Keq 1

37. Calculating Keq from E
Determine Keq for the following reaction.
Fe(s) + Pb2+(aq) → Pb(s) + Fe2+(aq) 1

38. Test Your Skill Determine Keq for the reaction below at 25 C.
2Ag(s) + Ni2+(aq) → Ni(s) + 2Ag+(aq) 1

39. The Nernst Equation
The Nernst equation is used to calculate cell potentials under non-standard conditions:
where Q is the reaction quotient. The second expression is for 25 C only. 1

40. The Nernst Equation
What is the cell potential for a cell composed of Zn, M Zn(NO3)2, Cu, and M Cu(NO3)2?
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) E = 1.10 V 1

41. Test Your Skill
What is the cell potential for a cell composed of Cu, M Cu(NO3)2, Fe, and M FeSO4? 1

42. Applications of Voltaic Cells
Measurement of species concentration: the pH meter.
The half-cell potential depends on the concentration of hydrogen ions in solution.
A pH electrode is an electrochemical cell with a known reference cell as a part.
A pH meter is simply a volt meter calibrated to display pH instead of volts. 1

43. Applications: Batteries
All batteries are voltaic cells.
In a dry cell, a Zn case is the anode, which is in contact with a moist paste of MnO2, NH4Cl, and a carbon electrode.
Cell reaction:
Zn(s) + 2NH4+(aq) + MnO2(s) →
Zn2+(aq) + Mn2O3(s) + 2NH3(aq) + H2O(l) 1

44. Dry Cell 1

45. Alkaline Dry Cell NH4Cl is replaced with KOH or NaOH.
Zn anode corrodes less.
Voltage is more constant.
Overall cell reaction:
Zn(s) + MnO2(s) → ZnO(s) + Mn2O3(s) 1

46. Alkaline Dry Cell 1

47. Lead Storage Battery Used for high current applications.
Overall reaction:
Pb(s) + PbO2(s) + 4H+ + 2SO42- →
2PbSO4(s) + 2H2O(l)
No salt bridge needed, because oxidizing and reducing agents do not come into contact with each other.
Rechargeable as long as the PbSO4 product adheres to electrodes (not limitless). 1

48. Lead Storage Battery 1

49. Fuel Cells
A fuel cell is a voltaic cell in which reactants are supplied continuously, products are removed continuously, and electrical energy is produced from the chemical reaction.
The H2/O2 fuel cell uses the reaction
2H2(g) + O2(g) → 2H2O(l)
and is used to produce electricity and water on space missions. 1