ELECTROCHEMISTRY(1).pptx

ELECTROCHEMISTRY
This is concerned with the transformation of chemical reactions into
electricity because energy is stored by chemical reagents
A science studying the relationship between chemical energy and electrical
energy and the rules of conversion of two energies.
Electrochemistry is the study of solutions of electrolytes and of phenomena
occurring at electrodes immersed in these solutions.
dmuloogi@must.ac.ug 1

Electric energy Chemical energy
Electrolysis
Galvanic cell

There are two kinds of substances which
conduct electricity
• Conductors- those in which charge is transferred by electrons e.g.
graphite and metals
• Electrolytes- those in which charge transfer is achieved by movement of
ions towards the oppositely charged electrodes.
• For conductors, their conductivity falls with rise in temperature because
the increase in temperature increases the vibration of atoms which
hinder the movement of electrons.
• In electrolytes, their conductivity increases with rise in temperature due
to decreased viscosity of water and increased velocity of the ions.

Difference between electronic & electrolytic
conductors
(3) Conduction increases with increase
in temperature
(3) Conduction decreases with increase
in temperature
(2) Flow of electricity is due to the
movement of ions
(2) Conduction is due to the flow of
electron
(1)Flow of electricity takes place by the
decomposition of the substance.
(1) Flow of electricity take place withou
t the decomposition of substance.
Electrolytic conductors
Electronic conductors

Conductivity and concentration

ELECTROLYTIC CONDUCTIVITY
• This is sometimes referred to as electrolytic conduction or specific
conductivity.
• Electrolytic conductivity is defined as the reciprocal of resistivity or
conductance of a solution enclosed between electrodes which are of unit
cross-sectional area (1 m2) and unit length apart (1 m).
• The resistance offered by solution to the flow of current is in accordance
with the one offered by metallic conductors in accordance with Ohm’s law.
At constant temperature, electric resistance (R) of an electrolyte is;
-Directly proportional to the distance (l) between the electrodes
-Inversely proportional to the cross-sectional area of the solution between
the electrodes (A)

Specific conductance
1
 

In summary
 
  
 
K
a
x Conductance
Unit of specific conductance is ohm–1cm–1
SI Unit of specific conductance is Sm–1 where S is Siemen
a
But ρ = R
K
a.R
 
l/a is known as cell constant
Conductance of unit volume of
cell is specific conductance.

DILUTION

Equivalent Conductance
Where, k = Specific conductivity
V = Volume of solution in cc. containing one
gram equivalent of the electrolyte.
It is the conductance of one gram equiv
alent of the electrolyte dissolved in V cc
of the solution.
Equivalent conductance is represented
by 
Mathematically,   
k V
1000
k
Normality
  

Measurement of conductivity of Solutions
• The conductivity of a solution is obtained by measuring its resistance
using a modified Wheatstone bridge circuit or a potentiometer

Variation of conductivity with concentration
(a) For strong electrolytes, conductivity depends on two factors;
(i) Concentration- conductivity increases with increase in the number of conducting ions i.e as concentration of ions
increases. However, at higher concentration, the number of ions per unit volume increases so that interionic interferences
increase due to ion-ion interaction, where at one point an ion (say cation) is surrounded by an atmosphere of oppositely
charged ions (say anions). This leads to formation of an ionic atmosphere which reduces the speed/velocity of the central
ion and hence decrease in conductivity.
(ii) Temperature- increase in temperature increases conductivity. This is because increase in temperature increases the
speed of conducting ions
(b) For weak electrolytes, conductivity depends on the degree of ionisation which increases with dilution (or reduces
with increase in concentration)

Variation of molar conductivity with
concentration.

WEAK ELECTROLYTES
Trend: Molar conductivity increases with decrease in concentration (or
decreases with increase in concentration)
Explanation: Molar conductivity increases with decrease in
concentration. At low concentration, more molecules ionise. The
degree of ionisation increases with decrease in concentration and thus
the number of ions per unit volume increases. Therefore, molar
conductivity also increases.

concentration, (mole L )
–1 1/2
CH COOH (weak electrolyte)
3
KCl (strong electrolyte)

Molar conductivity against dilution (1/C)
Strong electrolytes
Trend: 𝛬 increases with dilution (i.e increases with decrease in concentration)
Explanation: As dilution increases, 𝛬 increases. This is because, ions of
opposite charge become far apart when concentration reduces. Interionic
interference decreases leading to an increase in ionic mobility.
NB: At a certain concentration, all the ions are completely separated, so
increase in dilution further has no effect on molar conductivity, thus the
graph levels off. When extrapolated backwards, the molar conductivity
obtained is molar conductivity at infinite dilution.

Weak electrolytes
Trend: Molar conductivity increases with dilution
Explanation: the degree of ionisation of a weak electrolyte increases with dilution ( or increases with decrease in
concentration). This is because at low concentration, the number of conducting ions is high since the degree of ionisation for
weak electrolytes increases with dilution
The removal of H+ ions from the acid equilibrium by water disturbs the equilibrium. This shifts the equilibrium position
from the left to the right, and more acid ionises which increases the number of conducting ions.
The number of ions produced by a weak electrolyte depends on water added. This number increases indefinitely and
molar conductivity has no limit. This means that, molar conductivity at infinite dilution for a weak electrolyte cannot be
determined using graphs.
Application of graphs:
Used to obtain molar conductivity at infinite dilution for strong electrolytes.

Factors which affect the magnitude of Molar conductivity:
1. Concentration- molar conductivity reduces with increase in concentration.
2. The charge on the ion- Molar conductivity increases with increase in ionic charge. The greater the
ionic charge, the more strongly attracted is the ion, the greater is its mobility and greater the molar
conductivity.
3. Ionic radius- Molar conductivity reduces as ionic radius reduces. This is because, the smaller the
ionic radius, the greater the shell of water molecules attracted. This lowers the mobility and lowers
molar conductivity.
4. Temperature- Molar conductivity increases with increase in temperature. Increase in temperature
increases the velocity/speed of conducting ions, hence increases molar conductivity

Law of the independent migration of ions
Kohlrausch discovered relations between the values of
for different electrolytes. For example:
m


2 -
1
(
K
C
l
)
0
.
0
1
4
9
9
S
m
m
o
l
m


 
2 -
1
(
L
i
C
l
)
0
.
0
1
1
5
0
S
m
m
o
l
m


 
2 -
1
3
(
K
N
O
)
0
.
0
1
4
5
0
S
m
m
o
l
m


 
2 -
1
3
(
L
i
N
O
)
0
.
0
1
1
0
1
S
m
m
o
l
m


 

The difference in four pairs of salts having comm
on ion is always approximately constant.

2
-
1
3 3
(
K
C
l
)
(
L
i
C
l
)
(
K
N
O
)
(
L
i
N
O
)
0
.
0
0
3
4
9
S
m
m
o
l
m
m
m
m














2
-
1
3 3
(
K
C
l
)
(
K
N
O
)
(
L
i
C
l
)
(
L
i
N
O
)
0
.
0
0
0
4
9
S
m
m
o
l
m
m
m
m














This behavior indicates that ions in an extremely d
ilute solution migrate independently. There is no in
teraction between different ions. Therefore
, ,
m m m
 
  
   






For example
At 25℃，
(NaAc) = 91.0×10-4 S·m2·mol–1，
(HCl)=426.2×10-4 S·m2·mol–1，
(NaCl)=126.5×10-4 S·m2·mol–1，
What is the molar conductivity of HAc at 25℃?
m
 
m
 
m
 

+
m m m
(
N
a
A
c
)
=
(
N
a
)
+
(
A
c
)


  

+
m m m
(
H
C
l
)
=
(
H
)
+
(
C
l
)


  

+
m m m
(
N
a
C
l
)
=
(
N
a
)
+
(
C
l
)

  

+
m m m
(
H
A
c
)
=
(
H
)
+
(
A
c
)


  

+ +
m m m m
+
m m
=
(
H
) (
C
l
) (
N
a
) (
A
c
)
(
N
a
) (
C
l
)
   
 
 
  
 

  
 
m m m
=
(
H
C
l
) (
N
a
A
c
)(
N
a
C
l
)


  
 
=(426.3 +91.0–126.5)×10–4 S·m2·mol–1
=390.7×10–4 S·m2·mol–1
Solution

Applications/uses of conductivity
measurements
• Conductimetric titration
• Determination of formulae of complexes
• Determination of solubility product for sparingly soluble salts
• Determination of molar conductivity at infinite dilution of strong
electrolytes
• Determination of degree of ionisation of weak electrolytes.

Correction for Non-ideality of electrolyte
• Ionic activity or effective concentration, a is defined as a measure of
how much electrolyte in solution is actually participating in the
conductivity. It is a factor of actual concentration of the electrolyte i.e
a = ȣC;
Where ȣ is activity coefficient which is a correction factor and
dimensionless; 0 ≤ ȣ ≤ 1
As ȣ 1, a C
At infinite dilution, a = C

The Debye- Hückel Theory

Transport or Transference Numbers

Transport Numbers
• Anything that affects ionic mobility affects transport numbers
• This means that transport numbers are always accompanied by;
-concentration
-temperature
-Nature of electrolyte

Methods of determining transport numbers
• Moving boundary method
• Electrochemical cells method
• Hittorf method

Solubility Product
• The solubility product of an electrolyte is the product of the
concentrations of the ions in a saturated solution of a sparingly
soluble ionic compound raised to appropriate stoichiometric powers
i.e. the coefficients on the respective ions produced by the
electrolyte.

SOLUBILITY PRODUCT
When a sparingly soluble salt such as AgCl is shaken with water, a saturated solution is formed.
If the salt is a strong electrolyte, the dissolved part is completely ionised. There is thus an
equilibrium established between the unionised solid and the ions produced.
Thus, the solubility equation is written as;
MX(s) M+(aq) + X-(aq)
By the law of mass action; 𝐾𝑠𝑝 = 𝑀+
𝑋−
. This is called the solubility product expression
for the sparingly soluble salt MX.

NOTE:
(a) A solution of MX is unsaturated if 𝐾𝑠𝑝 > 𝑀+
𝑋−
. This solution can dissolve more
of the salt MX.
(b) The solution of MX is saturated if 𝐾𝑠𝑝 = 𝑀+
𝑋−
. This solution cannot dissolve
anymore salt.
(c) If 𝐾𝑠𝑝 < 𝑀+
𝑋−
, the salt MX will be precipitated.
(d) Ionic concentration can be increased by evaporation which deposits solid MX or by
adding a solution containing one of the ions of the electrolyte (common ionic effect).
The main use of solubility product is to predict whether an electrolyte will dissolve or
precipitate out of solution.
Examples of sparingly soluble electrolyte/salt include; AgCl, Ag2CO3, Ca(OH)2, PbCl2,
Ag2BrO3, BaSO4, Ca3(PO4)2.
NB: if on dissolution, the salt gives more than 1 mole of a certain kind of ion, the concentration
of that ion is raised to the corresponding power;
e.g. Lead(II) chloride has solubility equation;
PbCl2(s) Pb2+(aq) + 2Cl-(aq)
The solublity product expression is given by;
𝐾𝑠𝑝 = 𝑃𝑏2+
𝐶𝑙− 2
Note that, [Pb2+
] = ½[Cl-
].

Factors affecting the solubility of sparingly soluble salt
1. Common ion effect.
Common ion effect is defined as the suppression of the degree of dissociation of a
sparingly soluble electrolyte containing a common ion.
Question 1:
Consider the equilibrium of a sparingly soluble salt BaSO4;
BaSO4(s) Ba2+(aq) + SO4
2-(aq)
(i) State the effect of adding dilute sulphuric acid (or any soluble sulphate). Give
reasons.
Addition of dilute sulphuric acid to the above equilibrium produces H+
ions and
sulphate ions. The sulphate ion from sulphuric acid is said to be a common ion to
sulphate ion from BaSO4.

Applications of solubility product
• Purification of common salt
• Salting out of soap
• Softening of hard water using sodium carbonate

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