Oxidation & Reduction
Electrochemistry
BLB 10th Chapters 4, 20
Lab Summary
 Oxidation and Reduction (redox)
 Activity series with three metals (Zn, Mg, Cu)
 Writing equations for reactions
 Electrochemical cells – standard & non-standard
(Assemble for E° > 0.)
 Writing equations for cells (Reactant metal is the
most reactive; product metal the least.)
 E° calculation (cathode – anode)
21.1, 4.4 Oxidation-Reduction Reactions
 Oxidation
 Loss of electrons
 Increase in oxidation number
 Gain of oxygen or loss of hydrogen
 Reduction
 Gain of electrons
 Decrease in oxidation number
 Loss of oxygen or gain of hydrogen
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
 Oxidizing agent or oxidant – reactant that
contains the element being reduced; is itself
reduced
 Reducing agent or reductant – reactant that
contains the element being oxidized; is itself
oxidized
Redox Reactions
 Combustion, corrosion, metal production,
bleaching, digestion, electrolysis
 Metal oxidation



Activity Series (Table 4.5, p. 143)
Some metals are more easily oxidized and
form compounds than other metals.
Displacement reaction – metal or metal ion is
replaced through oxidation
A + BX → AX + B
20.3 Voltaic Cells
 A spontaneous redox reaction can perform
electrical work.
 The half-reactions must be placed in separate
containers, but connected externally.
 This creates a potential for electrons to flow.
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Line notation:
Zn(s)|Zn2+(aq)||Cu2+(aq)|Cu(s)
Movement of Electrons
e¯
Zn(s) → Zn2+(aq) + 2 e¯
Cu2+(aq) + 2 e¯ → Cu(s)
Net reaction: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
20.4 Cell EMF
 EMF – electromotive force – the potential energy
difference between the two electrodes of a
voltaic cell; Ecell; measured in volts
 E°cell – standard cell potential (or standard emf)
 For the Zn/Cu cell, E°cell = 1.10 V
 electrical work = Coulombs x volts
J=CxV
J
V
C
Standard Reduction (Half-cell) Potentials
 E° - potential of each half-cell
 E°cell = E°cell(cathode) - E°cell(anode)
 For a product-favored reaction:
 ΔG°
<0
 E°cell > 0
 Measured against standard hydrogen
electrode (SHE); assigned E° = 0 V.
Problem 28 Sketch and calculate the E°.
Voltaic cell with: Al(s) in Al(NO3)3(aq) on one
side and a SHE on the other.
20.5 Spontaneity of Redox Reactions
n = # moles of e¯ transferred
ΔG° = wmax = −nFE° F = 96,485 C/mol (Faraday constant)
wmax = max. work
 ΔG° < 0
 E°cell > 0
ΔG° for problem 28
20.6 Effect of Concentration on Cell EMF
 Concentrations change as a cell runs.
 When E = 0, the cell is dead and reaches equilibrium.
 Nernst equation allows us to calculate E under
nonstandard conditions:
J
R

8
.
3145
RT


mol K
E  E  
 ln Q
C
F  96,485 mol
 nF 
@ 298 K
 0.0257 
E  E  
 ln Q or
 n 
 0.0592 
E  E  
 log Q
 n 
Cell EMF and Equilibrium
 When E = 0, no net change in flow of electrons and
cell reaches equilibrium.
0.0257
E 
ln K
n
 and
nE 
ln K 
0.0257
K of problem 28
or
or
0.0592
E 
log K
n
nE 
log K 
0.0592