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Electrochemistry -1
Puneet Jyoti
G.C.G-42 Chd.
Electron Transfer Reactions
Results in the generation of an electric current (electricity) or be caused by
imposing an electric current.
Electron transfer reactions are oxidation-reduction or redox reactions.
Therefore, this field is often called ELECTROCHEMISTRY.
Electrochemistry is a branch of chemistry that studies chemical reactions which
take place in a solution at the interface of an electron conductor (a metal or a
semiconductors) and an ionic conductor (the electolyte), and which involve electron
transfer between the electrode and the electrolyte or species in solution.
Motion of ions in the solution:
1) diffusion: due to difference in
concentration
2) convection: due to the difference in
density or temperature
3) transfer: due to electric field
Only the transfer can cause net electricity
Conductance and its measurement
For metals:
Ohm’s Law
U
R
I
R: resistance
Dimension: Ohm, 
Resistivity
l
R
A
Dimension: Ohm m,  m
For electrolytic solution:
electric conductance (G) :
Definition: G = 1/R
Dimension: -1, mho, Siemens, S
conductivity () or spedific conductance:
Definition:  = 1/ 
Dimension: S m-1
Cell constant of a conductivity
cell
l 
  G   K cellG
 A
K cell  R
Influential factors for conductivity
1) concentration
– dependence of
conductance
1. Acids and bases
have higher
conductance
2. C < 5 mol dm-3,
 increases with C
3. For CH3COOH conductance does
not depend on C
(2) temperaturedependence of
conductance
Molar conductivity
The conductivity of a solution is approximately
proportional to the concentration
1) Definition
m 

C
m 

1
V
 V
V: degree of dilution
m is the conductivity contributed by 1 mole of
electrolyte between electrodes of 1 m apart
Dependence of molar conductivity
on concentration
m decreases with
concentration.
Due to the interaction
between ions:
interionic attraction
Kohlrausch replotted
m against C1/2
For 1:1
electrolytes:
C < 0.002~ 0.003
mol dm-3
Linear
relationship
between m and
C1/2 can be
observed.
Kohlrausch empirical formula

m
m    A c
To extrapolate the linear part of m ~ C1/2 at
low concentration to C = 0, m can be
obtained.
m the limiting value of m at infinite
dilution: limiting molar conductivity
Kohlrausch’s law of independent
ionic mobilities
 

m

m, 


m, 
At infinite dilution, m should be
the
sum
of
the
separate
contributions of the ions

limiting molar conductivity of
weak electrolyte
 (HAc)   (H )   ( Ac )

m

m


m

  ( H )   (Cl )   ( Na ) 

m


m


m

 ( Ac )   ( Na )   (Cl )

m


m

m


m

m


m
  (HCl)   ( NaAc)   ( NaCl)
Ionic mobility and transference
number of ions
1) Ionic mobility
dE
r
dl
dE
r U
dl
Under unit potential gradient: dE/dl = 1
V m-1: U = R, ionic mobility
2) Transference number
I = I + + IQ = Q + + Q-
tj 
Qj
Q
The fraction of the current transported by an ion
is its transference number or transport number
t = t + + t- = 1
3) Relation between ionic mobility and
transference number
C-, Z-, U-; C+, Z+, U+;
For time t:
Q+ = A U+t C+ Z+ F
Q  = A Ut C Z F
Q = Q+ + Q = AtF ( U+C+ Z+ + U C Z)
C+ Z+ = C Z
Q = AtF C+ Z+ ( U+ + U)
U
U
t 
t 
U U
U U
Measurement of transference numbers
1) Hittorf method (1853)
Electrolysis of HCl solution
Anodic region
Bulk solution
cathodic region
When 4 Faraday pass through the electrolytic cell
4 Cl- -4e-  2 Cl2
3 mol H+
1 mol Cl- 
4 H+ +4e-  2 H2
3 mol H+
1 mol Cl- 
For anodic region:
Cresidual = Cinitial – Creact + C transfer
3
=
6
–
t- = 1 / 4 = 0.25
4
+ C transfer
t+ = 3 / 4 = 0.75
Hittorf’s
transference cell
2) The moving-boundary
method
MA, MA’ have an ion in
common. The boundary,
rather difference in color,
refractivity, etc. is sharp.
In the steady state, the two
ions move with the same
velocity.
When Q coulomb passes, the
boundary moves x, the crosssectional area of the tube is
A:
xACZ+F = t+Q
THANX