CHAPTER 16 (pages 776-792)
1. Oxidation and Reduction
2. Galvanic Cells, Half Reactions (E°anode &
E°cathode)
3. Standard Reduction Potential (E°)
4. Nernst Equation, and the dependence of
Potential on Concentration
5. Relationship between Equilibrium
Constant and Standard Potential
6. Driving Force, ΔG and ε
1
REDOX REACTIONS
MnO2 + 4 HBr  MnBr2 + Br2 + 2 H2O
3 H2S + 2 NO3– + 2 H+  3 S + 2 NO + 4 H2O
2
OBSERVED REDOX PROCESSES
3
GALVANIC CELLS
4
INERT ELECTRODES
6
STANDARD REDUCTION POTENTIALS
7
8
MEASURING STANDARD POTENTIALS
9
CALCULATING STANDARD CELL
POTENTIAL
Al(s) + NO3−(aq) + 4 H+(aq)

Al3+(aq) + NO(g) + 2 H2O(l)
10
ADDITIONAL EXAMPLE
Fe(s) + Mg2+(aq)

Fe2+(aq) + Mg(s)
11
ox:
Fe(s)  Fe2+(aq) + 2 e−
red: Pb2+(aq) + 2 e−  Pb(s)
E = +0.45 V
E = −0.13 V
tot: Pb2+(aq) + Fe(s)  Fe2+(aq) + Pb(s)
E = +0.32 V
ELECTROMOTIVE POTENTIAL
13
E°CELL, G° AND K
Under standard state conditions, a reaction
will spontaneously proceeds in the forward
direction if:
– G° < 1 (negative)
– E° > 1 (positive)
–K>1
Design a voltaic cell with the following half
cells and complete the calculations:
Ag+ (aq) + 1e-  Ag (s) Eo = 0.80 V
Pb2+ (aq) + 2e-  Pb (s) Eo = -0.13 V
a. Calculate the Eocell (potential at standard
conditions)
b. Calculate Go.
c. Calculate 
d. Calculate the Ecell if [Ag+] = 2.0 M and
[Pb2+] = 1.0 x 10-4 M.
Williams, spring 2009
stop here
Design a voltaic cell with the following half
cells and complete the calculations:
Ag+ (aq) + 1e-  Ag (s) Eo = 0.80 V
Pb2+ (aq) + 2e-  Pb (s) Eo = -0.13 V
Calculate the Eocell (potential at standard
conditions)
Design a voltaic cell with the following half
cells and complete the calculations:
Ag+ (aq) + 1e-  Ag (s) Eo = 0.80 V
Pb2+ (aq) + 2e-  Pb (s) Eo = -0.13 V
Calculate Go.
Design a voltaic cell with the following half
cells and complete the calculations:
Ag+ (aq) + 1e-  Ag (s) Eo = 0.80 V
Pb2+ (aq) + 2e-  Pb (s) Eo = -0.13 V
Calculate 
Design a voltaic cell with the following half
cells and complete the calculations:
Ag+ (aq) + 1e-  Ag (s) Eo = 0.80 V
Pb2+ (aq) + 2e-  Pb (s) Eo = -0.13 V
Calculate the Ecell if [Ag+] = 2.0 M and Pb2+] = 1.0 x
10-4 M.
OBJECTIVE 11.4: PROVIDE A THOROUGH OVERVIEW OF
APPLICATIONS OF ELECTROCHEMICAL CELLS INCLUDING FUEL CELLS,
CORROSION, AND OTHER TOPICS AS TIME PERMITS.
23
CORROSION
• corrosion is the spontaneous oxidation of a
metal by chemicals in the environment
• since many materials we use are active
metals, corrosion can be a very big problem
RUSTING
•
•
•
•
rust is hydrated iron(III) oxide
moisture must be present
electrolytes promote rusting
acids promote rusting
– lower pH = lower E°red
Dry Cell Batteries
Lead – Acid Storage Battery
Biological Electrochemistry
Lithium Ion Battery