Oxidation-Reduction Reaction
※ Electrochemical reaction:
 The chemical reaction involved in electron transfer.
 Oxidation-reduction reaction.
※ Oxidation-reduction:
☆ Oxidation - the loss of electron, e.g. Zn  Zn  2e .
2

☆ Reduction - the gain of electron, e.g. Cu  2e  Cu .
2

☆ Oxidant - the species as an electron acceptor, e.g. Cu.
☆ Reluctant- the species as an electron donor, e.g. Zn.
☆ An electrochemical reaction - one oxidation and one
reduction.
☆ Oxidation - anode.
☆ Reduction - cathode.
※ Notation in electrochemical reaction:
 Two electrodes:
1
anode/cathode.
 Electrolyte solution.
 Salt bridge (not necessary).
☆Notation: Oxidation Reduction or Oxidation Reduction
 with liquid junction (salt bridge): A A
e.g. CuCu

 0.01m 
Ag

 0.02m 

B
 xm 

 ym 
B
Ag
 without liquid junction (salt bridge): A C
common ion in A and B); g. Pt H
2( g )

HCl
 xm 
B (C contains
(1 m )
AgCl Ag
( s)
※ Types of electrochemical cell:
★ Voltaic (Galvanic) cell:
 battery that store electrical energy.
 spontaneous reaction.
 electrons flow from the anode to the cathode via an
external conductor.
★ Electrolytic cell:
 requiring an external source of electrical energy for
operation.
 non-spontaneous reaction.
2
Anode
Cathode
Voltaic
-
+
Electrolytic
+
-
※ Cell potential:
 standard condition @P=1bar , T=25℃and C=1m.
 Ecell measured in reduction potential.
E  E
cell
anode
 Ecell > 0
E

cathode
.
voltaic cell.
※ Standard electrodes:
 Hydrogen electrode: Pt H ( 1bar ) H ( 1m )

( aq )
2
 Calomel electrode: Hg Hg Cl
2
3

2( S )
E0=0 V.
Cl ( 0 .1m ) E0=0.3338 V.
( aq )
※ Thermodynamics of electrochemical cells:
☆ Gibbs energy:
G  w
0
non  PV
 G   zFE for any standard reaction.
0
0
where F=96500 C/mole and z =charge number for cell
reaction.
 E0 > 0, spontaneous reaction.
E 
0
RT
n K
0
zF
☆ Electromotive force in non-standard condition:
For aA  bB  yY  zZ reaction:
EE
0
Y   Z 
 n
zF
 A  B
y
z
a
b
RT
@25℃, E  E 
0
0.0592
z
called the Nernst equation.
Y   Z 
og
 A  B
y
z
a
b
☆ Entropy:
 G 
 S   
 T 
P
 E 
S  zF  
 T 
 G 
 S  

 T 
P
P
☆ Enthalpy: H  G  T S
4
E 

H   zF  E  T 

T 
※ Nernst potentials:
☆ Situation I: K+ and Cl- are both permeable
 At equilibrium:
[Cl-] 1=[Cl-] 2; [K+] 1=[K+] 2
   0
☆ Situation II: K+ is only permeable
 Some K+ ions across membrane
from left to right.
 A potential established to
prevent more K+ ions from
crossing.
 The potential is called the Nernst
potential.  
RT
F
n
C1
C2
 Concentration change on both
sides is very little.
※ Types of electrochemical cells:
(I) Chemical cells: the cell is involved in a net chemical change.
5
 Cells in which the chemical reaction involves the
electrodes:
2
EE 
2
0
e.g. Zn Zn Cu Cu
RT
F
 Zn 
n
Cu 
2
2
 Redox cells: both the oxidized and reduced species in
solution and their interconversion is effected
by an inert electrode.
e.g. Pt H H
2

2
EE 
3
Fe , Fe Pt
( 1m )
0
 Fe 
n
 Fe 
2
RT
3
F
(II) Concentration cells: the cell is involved in a change of
concentration.
e.g. Pt H HCl
(m1 )
2
HCl
E
H Pt
( m2 )
2
RT
n
F
m2
m1
※ Applications of EMF measurements:
(A) pH determination:

Pt H ( 1bar ) H ( a ) Hg Cl
H
2
oxidation:
1
reduction:
1
overall:
 pH 
2
2
 H e

2( g )

2( S )
Hg Cl ( 1m )
(S)

Hg Cl  e  Hg  Cl

2
1
H
2
H
2
2( g )
 1
0
Hg Cl  Hg  Cl  H

2
2
2(S)
E  0. 2 8 0 2
0 .0 5 9 1
E  0 . 2802V

(S)
2
@25℃.
6
(S)

(B) Determination of activity coefficient:
Pt H ( 1bar ) HCl
2
1
oxidation:
H
2
AgCl Ag Cl
( aq )
(S)
(S)
 H e

2( g )


reduction: AgCl  e  Ag  Cl

overall:
1
EE 
RT
0
H
2
H
AgCl  Ag  Cl  H

0
H+
  m;a

Cl 

(S)
(S)
n a a  E 
F
a

2( g )

(S)
(S)
2RT
Cl 
n m 
F
2RT
n 
F

  m

(C) Determinatio
n of
dissociation constant, HA:
Pt H ( 1bar ) HA ( m ), NaA ( m ) AgCl Ag NaCl ( m )
2
oxidation:
1
1
2
H
2
 H e

2( g )
7

(S)
(S)
3
reduction: AgCl  e  Ag  Cl

overall:
1
EE 
RT
0
2
H
E 
RT
F
E 
0
RT
H+
(E  E )
F
RT

E 
0
Cl 
A
-
n K 
a
 n[
mA
RT
RT
-
n[
n[
a a a a

]
a
a
n K
Cl -
A-
]   n[
 HA Cl
-
A
-
RT
2
( aq )
AgCl Ag
(S)
8
(S)
A-
HA
A-
n[
F
] n K
-
(D) Determination of solubility product:
Cl ( 1bar ) HCl
Cl -
a
m m
]
m
HA
H+
HA
F
F
mHA mCl
RT
F
Cl -
HA

(S)
(S)
a a
n[
]
a
F
0
AgCl  Ag  Cl  H
n a a
F
0

2( g )

(S)
(S)
 
]

Cl -
HA
A-
a
oxidation: Cl 

1
2
Cl
2( g )
e

reduction: AgCl  e  Ag  Cl


(S)
(S)
overall: AgCl  Ag  Cl
+

(S)
EE 
0
RT
F
n
1
a Ag a Cl

E 
0
RT
F

9
n a a
Ag

E 
0
Cl

RT
F
n K
SP