1. Predicting Spontaneous Reactions

2. Standard Electrode 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)
19.3

3. E0 is for the reaction as written
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
19.3

4. Reduction Potentials
The half-reaction with the more positive reduction potential ( higher on the chart) is reduced and the other half-reaction is forced to oxidize.
Example: What spontaneous reaction occurs if Cl2 and Br2 are added to a solution of Cl- and Br-?

5. Standard Reduction Potentials (in Volts), 25oC Eo
Reaction Eo
F2 + 2e- ---> 2F-
+2.87
Co3+ + e- ---> Co2+
+1.80
PbO2 + 4H+ + SO e- ---> PbSO4(s) + 2H2O
+1.69
MnO4- + 8H+ + 5e- ---> Mn2+ + 4H2O
+1.49
PbO2 + 4H+ + 2e- ---> Pb2+ + 2H2O
+1.46
Cl2 + 2e- ---> 2Cl-
+1.36
Cr2O H+ + 6e- ---> 2Cr3+ + 7H2O
+1.33
O2 + 4H+ + 4e- ---> 2H2O
+1.23
Br2 + 2e- ---> 2Br-
+1.07

6. Example Continued: Cl2 + 2 e-  2 Cl- + 1.36 V
Br2 + 2 e-  2 Br V
Cl2 is more positive so it is reduced and Br2 is oxidized.
Write Br2 to show oxidation reaction:
2 Br-  Br2 + 2e V
Cl2 + 2 e-  2 Cl V
Cl2 + 2 Br-  Br2 + 2 Cl = EoCell

7. Spontaneous Reactions
E0Cell = reduction + oxidation potential
A positive Eocell means the reaction is spontaneous in that direction. A negative Eocell means the reverse reaction is spontaneous.

8. Standard Electrode Potentials
E0 = 0.34 V
cell
E0 = Ecathode + Eanode
cell
Ecell = ECu /Cu + E H /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)
19.3

9. Cd is the stronger oxidizer
What is the standard emf of an electrochemical cell made of a Cd electrode in a 1.0 M Cd(NO3)2 solution and a Cr electrode in a 1.0 M Cr(NO3)3 solution?
Cd2+ (aq) + 2e Cd (s) E0 = V
Cd is the stronger oxidizer
Cd will oxidize Cr
Cr3+ (aq) + 3e Cr (s) E0 = V
Anode (oxidation):
Cr (s) Cr3+ (1 M) + 3e-
x 2
Cathode (reduction):
2e- + Cd2+ (1 M) Cd (s)
x 3
2Cr (s) + 3Cd2+ (1 M) Cd (s) + 2Cr3+ (1 M)
E0 = Ecathode + Eanode
cell
E0 = (+0.74)
cell
E0 = 0.34 V
cell
19.3

10. Free Energy Change & Eocell
When a reaction takes place in a voltaic cell it performs work
W = nFEcell
n = moles of e- transferred
F = Faradays constant (96,485 C/mole e-)
Ecell = cell voltage

11. Work and Ecell
The maximum amount of work that a system cam do is equal to the negative of the change in Gibbs Free energy.
-DG =welectric = nFEcell
DG = -nFEcell
DGo = -nFEocell (standard conditions)
If Ecell = 0, the system is in equilibrium

12. Concentration and Ecell
The cell potential gradually drops as the reactants are consumed.
DG = DGo + RT ln Q
R = gas constant
T = temp (kelvin)
Q = reaction quotient (original molarities)

13. Concentration and Ecell
Substituting for DG and DGo
-nFEcell = -nFEocell + RT ln Q nF
The Nernst Equation
Ecell = Eocell – RT ln Q
Ecell = Eocell V log Q (base 10 log) n

14. Spontaneity of Redox Reactions
19.4

15. What is the equilibrium constant for the following reaction at 250C
What is the equilibrium constant for the following reaction at 250C? Fe2+ (aq) + 2Ag (s) Fe (s) + 2Ag+ (aq) = V n
ln K
Ecell
Oxidation:
2Ag Ag+ + 2e-
n = 2
Reduction:
2e- + Fe Fe
E0 = EFe /Fe + EAg /Ag 2+ +
E0 =
E0 = V V
x n E0
cell
exp
K = V
x 2
-1.24 V
= exp
K = 1.23 x 10-42
19.4

16. Fe2+ (aq) + Cd (s) Fe (s) + Cd2+ (aq)
Will the following reaction occur spontaneously at 250C if [Fe2+] = 0.60 M and [Cd2+] = M?
Fe2+ (aq) + Cd (s) Fe (s) + Cd2+ (aq)
Oxidation:
Cd Cd2+ + 2e-
n = 2
Reduction:
2e- + Fe Fe
E0 = EFe /Fe + ECd /Cd 2+ 2+
E0 = (-0.40) - V n
ln Q E
E =
E0 = V - V 2 ln
-0.04 V
E =
0.010
0.60
E = 0.013
E > 0
Spontaneous
19.5

17. Charging a Battery
When you charge a battery, you are forcing the electrons backwards (from the + to the -). To do this, you will need a higher voltage backwards than forwards. This is why the ammeter in your car often goes slightly higher while your battery is charging, and then returns to normal.
In your car, the battery charger is called an alternator. If you have a dead battery, it could be the battery needs to be replaced OR the alternator is not charging the battery properly.

18. Batteries Dry cell Leclanché cell Zn (s) Zn2+ (aq) + 2e- Anode:
Cathode:
2NH4 (aq) + 2MnO2 (s) + 2e Mn2O3 (s) + 2NH3 (aq) + H2O (l) +
Zn (s) + 2NH4 (aq) + 2MnO2 (s) Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s)
19.6

19. Batteries Mercury Battery Anode:
Zn(Hg) + 2OH- (aq) ZnO (s) + H2O (l) + 2e-
Cathode:
HgO (s) + H2O (l) + 2e Hg (l) + 2OH- (aq)
Zn(Hg) + HgO (s) ZnO (s) + Hg (l)
19.6

20. Batteries Lead storage battery Anode:
Pb (s) + SO2- (aq) PbSO4 (s) + 2e- 4
Cathode:
PbO2 (s) + 4H+ (aq) + SO2- (aq) + 2e PbSO4 (s) + 2H2O (l) 4
Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO2- (aq) PbSO4 (s) + 2H2O (l) 4
19.6

21. Solid State Lithium Battery
Batteries
Solid State Lithium Battery
19.6

22. Batteries
A fuel cell is an electrochemical cell that requires a continuous supply of reactants to keep functioning
Anode:
2H2 (g) + 4OH- (aq) H2O (l) + 4e-
Cathode:
O2 (g) + 2H2O (l) + 4e OH- (aq)
2H2 (g) + O2 (g) H2O (l)
19.6

23. Corrosion
19.7

24. Cathodic Protection of an Iron Storage Tank
19.7

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

26. Electrolysis of Water
19.8

27. Chemistry In Action: Dental Filling Discomfort
Hg2 /Ag2Hg V 2+
Sn /Ag3Sn V 2+
Sn /Ag3Sn V 2+

28. Electrolysis and Mass Changes
charge (Coulombs) = current (Amperes) x time (sec)
1 mole e- = 96,500 C = 1 Faraday
1 amp = 1 Coulomb / sec
19.8

29. How much Ca will be produced in an electrolytic cell of molten CaCl2 if a current of A is passed through the cell for 1.5 hours?
Anode:
2Cl- (l) Cl2 (g) + 2e-
Cathode:
2 mole e- = 1 mole Ca
Ca2+ (l) + 2e Ca (s)
Ca2+ (l) + 2Cl- (l) Ca (s) + Cl2 (g)
mol Ca = 0.452 C s
x 1.5 hr x 3600 s hr
96,500 C
1 mol e- x
2 mol e-
1 mol Ca x
= mol Ca
= 0.50 g Ca
19.8