1. Chapter 19
Electrochemistry
Semester 1/2009
Ref:
19.2 Galvanic Cells
19.3 Standard Reduction Potentials
19.4 Spontaneity of Redox Reactions
19.5 The Effect of Concentration on Emf
19.8 Electrolysis

2. 19.2 Galvanic Cells anode cathode oxidation reduction spontaneous
redox reaction

3. Cell = half-cell + half – cell Oxidation Reduction Anode Cathode
In Galvanic cell…
Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)
Zn is oxidized to Zn2+ ion
Zn electrode is Anode (Reducing Agent)
Cu2+ is reduced to Cu
 Cu electrode is Cathode (Oxidizing Agent)

4. Galvanic Cells
The difference in electrical potential between the anode and cathode is called:
cell voltage
electromotive force (emf)
cell potential
Cell Diagram
Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)
[Cu2+] = 1 M & [Zn2+] = 1 M
Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
anode
cathode

5. Standard Electrode Potentials
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)
Anode (oxidation):
Zn (s) Zn2+ (1 M) + 2e-
Cathode (reduction):
2e- + 2H+ (1 M) H2 (1 atm)
Zn (s) + 2H+ (1 M) Zn2+ + H2 (1 atm)

6. 19.3 Standard Reduction Potentials
Standard reduction potential (E0) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm.
Reduction Reaction
2e- + 2H+ (1 M) H2 (1 atm)
E0 = 0 V
Standard hydrogen electrode (SHE)

7. Standard Electrode Potentials
E0 = 0.76 V
cell
Standard emf (E0 )
cell
E0 = Ecathode - Eanode
cell
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)
E0 = EH /H - EZn /Zn
cell + 2+ 2
0.76 V = 0 - EZn /Zn 2+
EZn /Zn = V 2+
Zn2+ (1 M) + 2e Zn E0 = V

8. Standard Electrode Potentials
E0 = 0.34 V
cell
E0 = Ecathode - Eanode
cell
Ecell = ECu /Cu – EH /H 2+ + 2
0.34 = ECu /Cu - 0 2+
ECu /Cu = 0.34 V 2+
Pt (s) | H2 (1 atm) | H+ (1 M) || Cu2+ (1 M) | Cu (s)
Anode (oxidation):
H2 (1 atm) H+ (1 M) + 2e-
Cathode (reduction):
2e- + Cu2+ (1 M) Cu (s)
H2 (1 atm) + Cu2+ (1 M) Cu (s) + 2H+ (1 M)

9. Note:
The more positive E0 the greater the tendency for the substance to be reduced
The half-cell reactions are reversible
The sign of E0 changes when the reaction is reversed
Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E0

10. 19.4 Spontaneity of Redox Reactions
DG = -nFEcell
n = number of moles of electrons in reaction
F = 96,500 J
V • mol
DG0 = -nFEcell
= 96,500 C/mol
DG0 = -RT ln K
= -nFEcell
Ecell = RT nF
ln K
(8.314 J/K•mol)(298 K)
n (96,500 J/V•mol)
ln K = = V n
ln K
Ecell = V n
log K
Ecell
E0 > 0 spontaneous reaction

11. Spontaneity of Redox Reactions

12. 19.5 The Effect of Concentration on Cell Emf
DG = DG0 + RT ln Q
DG = -nFE
-nFE = -nFE0 + RT ln Q
DG0 = -nFE
Nernst equation
E = E0 -
ln Q RT nF
At 298 K - V n
ln Q E
E = - V n
log Q E
E =

13. 19.8 Electrolysis is the process in which electrical energy is used to cause a nonspontaneous chemical reaction to occur.

14. Electrolysis of Water
19.8

15. Electrolysis and Mass Changes
Quantitative Aspects
Case (i) Na + + 1e Na
1 mol. of electron produces 1 mol of Na Atom
1 F (96500 C)
Case (ii) Mg e Mg
2 mol. of electron produces 1 mol of Mg Atom
2 F (2x 96500C)
Case (iii) Al e Al
3 mol. of electron produces 1 mol of Al Atom
3 F (3 x C)

16. charge ( C ) = current (A) x time (s)
1 mole of electron = coulomb
1 mol. of Na atom = 22 g
1 mol. of Mg atom = 24 g
1 mol. of Al atom = 26 g