The document covers redox reactions and electrochemistry, including topics such as oxidation-reduction reactions, galvanic cells, and standard reduction potentials. It explains the processes involved in balancing redox reactions, the operation of galvanic cells, and how to determine cell potential and free energy changes in these systems. Various examples illustrate the concepts, calculations of standard cell potentials, and the effects of concentration on cell potential.

1. Redox Reactions and Electrochemistry
• Redox reactions
• Galvanic cells
• Standard reduction potential
• Cell potential
– Free energy
– Effect of concentration
• Battery
• Corrosion
• Electrolysis

2. I. Oxidization–Reduction (Redox) Reaction
2Fe 2Fe2+ + 4e-
O2 + 4e- 2O2-
Oxidation half-reaction (lose e-):
Fe is oxidized and Fe is the reducing agent
Reduction half-reaction (gain e-):
O2 is reduced and O2 is the oxidizing agent
2Fe (s) + O2 (g) 2FeO (s)
Electrochemical processes are redox reactions in which:
• energy released by a spontaneous reaction is converted to electricity
• electrical energy is used to cause a nonspontaneous reaction
0 0 2+ 2-

3. A. Oxidation number/state
• Free elements have an oxidation number (ON) =0
• Monatomic ions, the ON is equal to the charge on the ion.
• The ON of oxygen is usually –2. In H2O2 and O2
2-:–1, in O2
- : -1/2
• The ON of hydrogen is +1 except when it is bonded to metals in
binary compounds. In these cases, its oxidation number is –1.
• Fluorine is always –1, Group IA metals are +1, IIA metals are +2
in compounds.
• The sum of the ON of all the atoms in a molecule or ion is equal to
the charge on the molecule or ion.
Ex. Oxidation numbers of all the atoms in HSO3
–
O =-2 H= +1 S=? 1 +3x (-2)+S= -1
S=4

4. B. Balance redox reaction
1. By inspection
2. Ion-electron method (half-reaction method)
By balancing the number of electrons lost and gained
during the reaction
Cu+ H+ Cu2+ + H2
2
Sn2+ (aq) + Fe3+ (aq) Sn4+(aq) + Fe2+(aq)
Oxidation half reaction
Reduction half reaction
Sn2+ (aq) Sn4+(aq) + 2 e–
Fe3+ (aq) + e– Fe2+(aq)
2 2
( ) x 2
2 Fe3+ (aq) + 2e– 2Fe2+(aq)

5. Ion-electron method
Fe2+ + Cr2O7
2- Fe3+ + Cr3+
2. Separate the equation into two half-reactions.
Oxidation:
Cr2O7
2- Cr3+
+6 +3
Reduction:
Fe2+ Fe3+
+2 +3
3. Balance the atoms other than O and H in each half-reaction.
Cr2O7
2- 2Cr3+
1. Write the unbalanced equation for the reaction ion ionic form.
The oxidation of Fe2+ to Fe3+ by Cr2O7
2- in acid solution?

6. Ion-electron method
4. For reactions in acid, add H2O to balance O atoms and H+ to
balance H atoms.
Cr2O7
2- 2Cr3+ + 7H2O
14H+ + Cr2O7
2- 2Cr3+ + 7H2O
5. Add electrons to one side of each half-reaction to balance the
charges on the half-reaction.
Fe2+ Fe3+ + 1e-
6e- + 14H+ + Cr2O7
2- 2Cr3+ + 7H2O
6. If necessary, equalize the number of electrons in the two half-
reactions by multiplying the half-reactions by appropriate
coefficients.
6Fe2+ 6Fe3+ + 6e-
6e- + 14H+ + Cr2O7
2- 2Cr3+ + 7H2O

7. Ion-electron method
7. Add the two half-reactions together and balance the final equation
by inspection. The number of electrons on both sides must
cancel.
6e- + 14H+ + Cr2O7
2- 2Cr3+ + 7H2O
6Fe2+ 6Fe3+ + 6e-
Oxidation:
Reduction:
14H+ + Cr2O7
2- + 6Fe2+ 6Fe3+ + 2Cr3+ + 7H2O
8. Verify that the number of atoms and the charges are balanced.
14x1 – 2 + 6x2 = 24 = 6x3 + 2x3
9. For reactions in basic solutions, add OH- to both sides of the
equation for every H+ that appears in the final equation.

8. Summary of balancing redox reaction in acidic solutions
1. Divide reaction into two incomplete half-reactions
2. Balance each half-reaction by doing the following:
a. Balance all elements, except O and H
b. Balance O by adding H2O
c. Balance H by adding H+
d. Balance charge by adding e– as needed
3. If the electrons in one half-reaction do not balance those in the
other, then multiply each half-reaction to get a common multiple
4. The overall reaction is the sum of the half-reactions
5. The oxidation half-reaction is the one that produces electrons as
products, and the reduction half-reaction is the one that uses
electrons as reactants

9. Summary of balancing redox reaction in basic solutions
Steps 1-3 are the same as the reaction in acidic solution
4. Since we are under basic conditions, we neutralize any H+ ions
in solution by adding an equal number of OH– to both sides
5. The overall reaction is the sum of the half-reactions
Ex. Balance I– + OCl– I2 + Cl– in basic solution?
Oxidation I– I2
Reduction
OCl– Cl–
2 H2O+2 e– + OCl– Cl– + H2O + 2 OH–
2 e– + H2O + OCl– Cl– + 2 OH–
2 I– I2 + 2 e–
2 I– + H2O + OCl– I2 + Cl– + 2 OH–
Add up
+ H2O
2 H+ +
2 OH– + + 2 OH–
2e– +

10. II. Galvanic Cells
• Consider reaction Zn(s) + Cu(NO3)2 --> Cu(s) + Zn(NO3)2
Spontaneous (activity series, chapter 4)
half reaction: Zn(s) --> Zn2+ + 2e– Cu2+ +2 e– --> Cu(s)
• Galvanic cell: a device uses the spontaneous redox reaction to
produce an electric current
Oxidation
Zn --> Zn2+ + 2e–
Anode
Reduction
Cu2+ +2 e– --> Cu
Cathode
e– flow
e–
Salt bridge
Flow of cation
Flow of anion
Wire
Electrode: anode (oxidation rxn) and cathode (reduction rxn)
Salt bridge: contains conc. soln of strong electrolyte (KCl)

11. II. Galvanic Cells
anode
oxidation
cathode
reduction

12. II. Galvanic cell
Cell diagram : a conventional notation to represent galvanic cell
Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)
[Cu2+] = 1 M and [Zn2+] = 1 M
Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
anode cathode
1. Anode is on the left, cathode on the right
2. | : represent boundary between phases
3. || : represent salt bridge
Ex. Draw cell diagram for the following reaction:
Al (s) + Ag+ (aq) Al3+ (s) + Ag (s)
Al (s) | Al3+ || Ag+ | Ag (s)

13. III. Standard reduction potential
A. Cell voltage, cell potential, electromotive force (emf)
1. The potential energy of the e– is higher at the anode than at the cathode
2. The difference in electrical potential between the anode and cathode is
called as emf
3. Unit: V (volt) = 1 J/C
C: coulomb, unit of charge
4. Standard cell potential, Eo
cell
Cell potential under standard condition
Zn(s)| Zn2+(1 M) || H+ (1 M)| H2 (1 atm)| Pt (s)
2e- + 2H+ (1 M) H2 (1 atm)
Zn (s) Zn2+ (1 M) + 2e-
Anode (oxidation):
Cathode (reduction):
Zn (s) + 2H+ (1 M) Zn2+ + H2 (1 atm)

14. B. Standard reduction potential
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.
E0 = 0 V
Standard hydrogen electrode (SHE) : half-cell contain 1M H+ and
1 atm H2, used as a reference to determine Eo for other species
2e- + 2H+ (1 M) H2 (1 atm)
Reduction Reaction

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

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

17. • 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

18. 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 = -0.40 V
Cr3+ (aq) + 3e- Cr (s) E0 = -0.74 V
Cd is the stronger oxidizer
Cd will oxidize Cr
2e- + Cd2+ (1 M) Cd (s)
Cr (s) Cr3+ (1 M) + 3e-
Anode (oxidation):
Cathode (reduction):
2Cr (s) + 3Cd2+ (1 M) 3Cd (s) + 2Cr3+ (1 M)
x 2
x 3
E0 = Ecathode - Eanode
cell
0 0
E0 = -0.40 – (-0.74)
cell
E0 = 0.34 V
cell
Ex.

19. Calculate Eo
cell for the following reactions:
(a) 2 Al (s) + 6 H+ -> 2 Al3+ + 3 H2 (g)
(b) H2O2 + 2 H+ + Cu (s) -> Cu2+ + 2 H2O
6 H+(aq) + 6e- 3 H2(g)
2 Al(s) 2 Al3+ + 6e-
Anode (oxidation):
Cathode (reduction):
Ex.
(a) 2 Al (s) + 6 H+ -> 2 Al3+ + 3 H2 (g)
E0 = EH /H - EAl /Al
cell
0 0
+ 3+
2 = 1.66 V
E0 = Ecathode - Eanode
cell
0 0
H2O2+2 H+ +2e- 2 H2O
Cu(s) Cu2+ + 2e-
Anode (oxidation):
Cathode (reduction):
(b) H2O2 + 2 H+ + Cu (s) -> Cu2+ + 2 H2O
= 1.77 V- 0.34 V = 1.43 V
E0 = E - ECu /Cu
cell
0 2+
H2O2/H2 O

20. IV. Cell Potential
• Free energy : spontaneity of redox reactions
DG = -nFEcell
DG0 = -nFEcell
0
n = number of moles of electrons in reaction
F = 96,500
J
V • mol
= 96,500 C/mol
DG0 = -RT ln K = -nFEcell
0
Ecell
0 =
RT
nF
ln K
(8.314 J/K•mol)(298 K)
n (96,500 J/V•mol)
ln K
=
=
0.0257 V
n
ln K
Ecell
0
=
0.0592 V
n
log K
Ecell
0

21. A. Free energy and spontaneity of redox reaction
DG = -nFEcell
DG0 = -nFEcell
0
=
0.0257 V
n
ln K
Ecell
0
=
0.0592 V
n
log K
Ecell
0

22. Ex.
2e- + Fe2+ Fe
2Ag 2Ag+ + 2e-
Oxidation:
Reduction:
=
0.0257 V
n
ln K
Ecell
0
E0 = -0.44 – (0.80) E0 = -1.24 V
0.0257 V
x n
E0
cell
exp
K =
n = 2
0.0257 V
x 2
-1.24 V
= exp
K = 1.23 x 10-42
E0 = EFe /Fe – EAg /Ag
0 0
2+ +
What is the equilibrium constant for the following reaction at
250C? Fe2+ (aq) + 2Ag (s) Fe (s) + 2Ag+ (aq)

23. B. Concentration effect on cell potential
DG = DG0 + RT ln Q DG = -nFE DG0 = -nFE0
-nFE = -nFE0 + RT ln Q
E = E0 - ln Q
RT
nF
Nernst equation
At 298 K
-
0.0257 V
n
ln Q
E0
E = -
0.0592 V
n
log Q
E0
E =

24. Ex.
2e- + Fe2+ Fe
Cd Cd2+ + 2e-
Oxidation:
Reduction:
n = 2
= -0.44 – (-0.40) V = -0.04 V
E0 = EFe /Fe – ECd /Cd
0 0
2+ 2+
-
0.0257 V
n
ln Q
E0
E =
-
0.0257 V
2
ln
-0.04 V
E =
0.010
0.60
E = 0.013 E > 0 Spontaneous
Will the following reaction occur spontaneously at 250C if
[Fe2+] = 0.60 M and [Cd2+] = 0.010 M?
Fe2+ (aq) + Cd (s) Fe (s) + Cd2+ (aq)
[Cd2+]
[Fe2+]
Q =

25. Ex.
2e- +2 Ag+ 2 Ag
Ni Ni2+ + 2e-
Oxidation:
Reduction:
n = 2
= 0.80 – (-0.25) V = 1.05V
E0 = EAg /Ag – E Ni /Ni
0 0
+ 2+
A redox reaction for the oxidation of nickel by silver ion is set
up with an initial concentration of 5.0 M silver ion and 0.050 M
nickel ion. What is the cell EMF at 298K?
-
0.0257 V
n
ln Q
E0
E =
-
0.0257 V
2
ln
1.05 V
E =
0.050
5.02 E = 1.13
[Ni2+]
[Ag+]2
Q =
Ni + 2 Ag+ Ni2+ + 2 Ag

26. V. Battery
Leclanché cell
Dry 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)
About 1.5 V

27. V. Battery
Zn(Hg) + 2OH- (aq) ZnO (s) + H2O (l) + 2e-
Anode:
Cathode: HgO (s) + H2O (l) + 2e- Hg (l) + 2OH- (aq)
Zn(Hg) + HgO (s) ZnO (s) + Hg (l)
Mercury Battery
1.35 V

28. V. Battery
Anode:
Cathode:
Lead storage
battery
PbO2 (s) + 4H+ (aq) + SO2- (aq) + 2e- PbSO4 (s) + 2H2O (l)
4
Pb (s) + SO2- (aq) PbSO4 (s) + 2e-
4
Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO2- (aq) 2PbSO4 (s) + 2H2O (l)
4
Total 12V

29. VI. Corrosion
• Deterioration of metal by an electrochemical process

30. VI. Corrosion
• Cathodic Protection of an Iron Storage Tank

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

32. A. Electrolysis of water

33. B. Quantitative aspect of electrolysis
charge (C) = current (A) x time (s)
1 mole e- = 96,500 C
Ex. It will take the greates amount electricity to produce 2 mol of
the ____ metal by electrolysis.
(1) potassium (2) calcium (3) aluminum (4) silver

34. Ex.
Anode:
Cathode: Ca2+ (l) + 2e- Ca (s)
2Cl- (l) Cl2 (g) + 2e-
Ca2+ (l) + 2Cl- (l) Ca (s) + Cl2 (g)
2 mole e- = 1 mole Ca
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
= 0.0126 mol Ca
= 0.50 g Ca
How much Ca will be produced in an electrolytic cell of molten
CaCl2 if a current of 0.452 A is passed through the cell for 1.5
hours?