NCERT Textbook: Electrochemistry | Chemistry Class 12 - NEET PDF Download

 Page 1


Electrochemistry is the study of production of
electricity from energy released during spontaneous
chemical reactions  and the use of electrical energy
to bring about non-spontaneous chemical
transformations. The subject is of importance both
for theoretical and practical considerations. A large
number of metals, sodium hydroxide, chlorine,
fluorine and many other chemicals are produced by
electrochemical methods. Batteries and fuel cells
convert chemical energy into electrical energy and are
used on a large scale in various instruments and
devices. The reactions carried out electrochemically
can be energy efficient and less polluting. Therefore,
study of electrochemistry is important for creating new
technologies that are ecofriendly. The transmission of
sensory signals through cells to brain and vice versa
and communication between the cells are known to
have electrochemical origin. Electrochemistry, is
therefore, a very vast and interdisciplinary subject. In
this Unit, we will cover only some of its important
elementary aspects.
After studying this Unit, you will be
able to
· describe an electrochemical cell
and differentiate between galvanic
and electrolytic cells;
· apply Nernst equation for
calculating the emf of galvanic cell
and define standard potential of
the cell;
· derive relation between standard
potential of the cell, Gibbs energy
of cell reaction and its equilibrium
constant;
· define resistivity (r), conductivity
(k) and molar conductivity (?
m
) of
ionic solutions;
· differentiate between ionic
(electrolytic) and electronic
conductivity;
· describe the method for
measurement of conductivity of
electrolytic solutions and
calculation of their molar
conductivity;
· justify the variation of
conductivity and molar
conductivity of solutions with
change in their concentration and
define ?
m
? (molar conductivity at
zero concentration or infinite
dilution);
· enunciate Kohlrausch law and
learn its applications;
· understand quantitative aspects
of electrolysis;
· describe the construction of some
primary and secondary batteries
and fuel cells;
· explain corrosion as an
electrochemical process.
Objectives
Chemical reactions can be used to produce electrical energy,
conversely, electrical energy can be used to carry out chemical
reactions that do not proceed spontaneously.
2
Electrochemistry
Unit Unit Unit Unit Unit
2
Electrochemistry
Reprint 2024-25
Page 2


Electrochemistry is the study of production of
electricity from energy released during spontaneous
chemical reactions  and the use of electrical energy
to bring about non-spontaneous chemical
transformations. The subject is of importance both
for theoretical and practical considerations. A large
number of metals, sodium hydroxide, chlorine,
fluorine and many other chemicals are produced by
electrochemical methods. Batteries and fuel cells
convert chemical energy into electrical energy and are
used on a large scale in various instruments and
devices. The reactions carried out electrochemically
can be energy efficient and less polluting. Therefore,
study of electrochemistry is important for creating new
technologies that are ecofriendly. The transmission of
sensory signals through cells to brain and vice versa
and communication between the cells are known to
have electrochemical origin. Electrochemistry, is
therefore, a very vast and interdisciplinary subject. In
this Unit, we will cover only some of its important
elementary aspects.
After studying this Unit, you will be
able to
· describe an electrochemical cell
and differentiate between galvanic
and electrolytic cells;
· apply Nernst equation for
calculating the emf of galvanic cell
and define standard potential of
the cell;
· derive relation between standard
potential of the cell, Gibbs energy
of cell reaction and its equilibrium
constant;
· define resistivity (r), conductivity
(k) and molar conductivity (?
m
) of
ionic solutions;
· differentiate between ionic
(electrolytic) and electronic
conductivity;
· describe the method for
measurement of conductivity of
electrolytic solutions and
calculation of their molar
conductivity;
· justify the variation of
conductivity and molar
conductivity of solutions with
change in their concentration and
define ?
m
? (molar conductivity at
zero concentration or infinite
dilution);
· enunciate Kohlrausch law and
learn its applications;
· understand quantitative aspects
of electrolysis;
· describe the construction of some
primary and secondary batteries
and fuel cells;
· explain corrosion as an
electrochemical process.
Objectives
Chemical reactions can be used to produce electrical energy,
conversely, electrical energy can be used to carry out chemical
reactions that do not proceed spontaneously.
2
Electrochemistry
Unit Unit Unit Unit Unit
2
Electrochemistry
Reprint 2024-25
32 Chemistry
Cu
E
ext
>1.1
e
–
Current
Cathode
+ve
Anode
–ve
Zn
Fig. 2.2
Functioning of Daniell
cell when external
voltage E
ext
 opposing the
cell potential is applied.
We had studied the construction and functioning of Daniell cell
(Fig. 2.1). This cell converts the chemical energy liberated during the
redox reaction
Zn(s) + Cu
2+
(aq) ® Zn
2+
(aq) + Cu(s) (2.1)
to electrical energy and has an electrical
potential equal to 1.1 V when concentration
of Zn
2+ 
and Cu
2+
 ions is unity (1 mol dm
–3
)
*
.
Such a device is called a galvanic or a
voltaic cell.
If an external opposite potential is applied
in the galvanic cell [Fig. 2.2(a)] and increased
slowly, we find that the reaction continues to
take place till the opposing voltage reaches
the value 1.1 V [Fig. 2.2(b)] when, the reaction
stops altogether and no current flows through
the cell. Any further increase in the external
potential again starts the reaction but in the
opposite direction [Fig. 2.2(c)]. It now functions
as an electrolytic cell, a device for using
electrical energy to carry non-spontaneous
chemical reactions. Both types of cells are
quite important and we shall study some of
their salient features in the following pages.
*Strictly speaking activity should be used instead of concentration. It is directly proportional to concentration. In dilute
solutions, it is equal to concentration. You will study more about it in higher classes.
2.1 2.1 2.1 2.1 2.1 Electrochemical Electrochemical Electrochemical Electrochemical Electrochemical
Cells Cells Cells Cells Cells
Fig. 2.1: Daniell cell having electrodes of zinc and
copper dipping in the solutions of their
respective salts.
salt
bridge
Zn Cu
anode
cathode current
ZnSO
4
CuSO
4
E <
ext
1.1V
e
-ve +ve
I=0
Zn Cu
ZnSO
4
CuSO
4
E =
ext
1.1V
When E
ext
 < 1.1 V
(i) Electrons flow from Zn rod to
Cu rod hence current flows
from Cu to Zn.
(ii) Zn dissolves at anode and
copper deposits at cathode.
When E
ext
 = 1.1 V
(i) No flow of
electrons or
current.
(ii) No chemical
reaction.
When E
ext
 > 1.1 V
(i) Electrons flow
from Cu to Zn
and current flows
from Zn to Cu.
(ii) Zinc is deposited
at the zinc
electrode and
copper dissolves at
copper electrode.
(a)
(b)
(c)
Reprint 2024-25
Page 3


Electrochemistry is the study of production of
electricity from energy released during spontaneous
chemical reactions  and the use of electrical energy
to bring about non-spontaneous chemical
transformations. The subject is of importance both
for theoretical and practical considerations. A large
number of metals, sodium hydroxide, chlorine,
fluorine and many other chemicals are produced by
electrochemical methods. Batteries and fuel cells
convert chemical energy into electrical energy and are
used on a large scale in various instruments and
devices. The reactions carried out electrochemically
can be energy efficient and less polluting. Therefore,
study of electrochemistry is important for creating new
technologies that are ecofriendly. The transmission of
sensory signals through cells to brain and vice versa
and communication between the cells are known to
have electrochemical origin. Electrochemistry, is
therefore, a very vast and interdisciplinary subject. In
this Unit, we will cover only some of its important
elementary aspects.
After studying this Unit, you will be
able to
· describe an electrochemical cell
and differentiate between galvanic
and electrolytic cells;
· apply Nernst equation for
calculating the emf of galvanic cell
and define standard potential of
the cell;
· derive relation between standard
potential of the cell, Gibbs energy
of cell reaction and its equilibrium
constant;
· define resistivity (r), conductivity
(k) and molar conductivity (?
m
) of
ionic solutions;
· differentiate between ionic
(electrolytic) and electronic
conductivity;
· describe the method for
measurement of conductivity of
electrolytic solutions and
calculation of their molar
conductivity;
· justify the variation of
conductivity and molar
conductivity of solutions with
change in their concentration and
define ?
m
? (molar conductivity at
zero concentration or infinite
dilution);
· enunciate Kohlrausch law and
learn its applications;
· understand quantitative aspects
of electrolysis;
· describe the construction of some
primary and secondary batteries
and fuel cells;
· explain corrosion as an
electrochemical process.
Objectives
Chemical reactions can be used to produce electrical energy,
conversely, electrical energy can be used to carry out chemical
reactions that do not proceed spontaneously.
2
Electrochemistry
Unit Unit Unit Unit Unit
2
Electrochemistry
Reprint 2024-25
32 Chemistry
Cu
E
ext
>1.1
e
–
Current
Cathode
+ve
Anode
–ve
Zn
Fig. 2.2
Functioning of Daniell
cell when external
voltage E
ext
 opposing the
cell potential is applied.
We had studied the construction and functioning of Daniell cell
(Fig. 2.1). This cell converts the chemical energy liberated during the
redox reaction
Zn(s) + Cu
2+
(aq) ® Zn
2+
(aq) + Cu(s) (2.1)
to electrical energy and has an electrical
potential equal to 1.1 V when concentration
of Zn
2+ 
and Cu
2+
 ions is unity (1 mol dm
–3
)
*
.
Such a device is called a galvanic or a
voltaic cell.
If an external opposite potential is applied
in the galvanic cell [Fig. 2.2(a)] and increased
slowly, we find that the reaction continues to
take place till the opposing voltage reaches
the value 1.1 V [Fig. 2.2(b)] when, the reaction
stops altogether and no current flows through
the cell. Any further increase in the external
potential again starts the reaction but in the
opposite direction [Fig. 2.2(c)]. It now functions
as an electrolytic cell, a device for using
electrical energy to carry non-spontaneous
chemical reactions. Both types of cells are
quite important and we shall study some of
their salient features in the following pages.
*Strictly speaking activity should be used instead of concentration. It is directly proportional to concentration. In dilute
solutions, it is equal to concentration. You will study more about it in higher classes.
2.1 2.1 2.1 2.1 2.1 Electrochemical Electrochemical Electrochemical Electrochemical Electrochemical
Cells Cells Cells Cells Cells
Fig. 2.1: Daniell cell having electrodes of zinc and
copper dipping in the solutions of their
respective salts.
salt
bridge
Zn Cu
anode
cathode current
ZnSO
4
CuSO
4
E <
ext
1.1V
e
-ve +ve
I=0
Zn Cu
ZnSO
4
CuSO
4
E =
ext
1.1V
When E
ext
 < 1.1 V
(i) Electrons flow from Zn rod to
Cu rod hence current flows
from Cu to Zn.
(ii) Zn dissolves at anode and
copper deposits at cathode.
When E
ext
 = 1.1 V
(i) No flow of
electrons or
current.
(ii) No chemical
reaction.
When E
ext
 > 1.1 V
(i) Electrons flow
from Cu to Zn
and current flows
from Zn to Cu.
(ii) Zinc is deposited
at the zinc
electrode and
copper dissolves at
copper electrode.
(a)
(b)
(c)
Reprint 2024-25
33 Electrochemistry
As mentioned earlier a galvanic cell is an electrochemical cell that
converts the chemical energy of a spontaneous redox reaction into
electrical energy. In this device the Gibbs energy of the spontaneous
redox reaction is converted into electrical work which may be used for
running a motor or other electrical gadgets  like heater, fan,
geyser, etc.
Daniell cell discussed earlier is one such cell in which the following
redox reaction occurs.
Zn(s) + Cu
2+
(aq) ® Zn
2+
 (aq) + Cu(s)
This reaction is a combination of two half reactions whose addition
gives the overall cell reaction:
(i) Cu
2+
   +  2e
–
    ® Cu(s) (reduction half reaction) (2.2)
(ii) Zn(s)  ® Zn
2+
 + 2e
–
(oxidation half reaction) (2.3)
These reactions occur in two different portions of the Daniell cell.
The reduction half reaction occurs on the copper electrode while the
oxidation half reaction occurs on the zinc electrode. These two portions
of the cell are also called half-cells or redox couples. The copper
electrode may be called the reduction half cell and the zinc electrode,
the oxidation half-cell.
We can construct innumerable number of galvanic cells on the pattern
of Daniell cell by taking combinations of different half-cells. Each half-
cell consists of a metallic electrode dipped into an electrolyte. The two
half-cells are connected by a metallic wire through a voltmeter and a
switch externally. The electrolytes of the two half-cells are connected
internally through a salt bridge as shown in Fig. 2.1.  Sometimes, both
the electrodes dip in the same electrolyte solution and in such cases we
do not require a salt bridge.
At each electrode-electrolyte interface there is a tendency of metal
ions from the solution to deposit on the metal electrode trying to make
it positively charged. At the same time, metal atoms of the electrode
have a tendency to go into the solution as ions and leave behind the
electrons at the electrode trying to make it negatively charged. At
equilibrium, there is a separation of charges and depending on the
tendencies of the two opposing reactions, the electrode may be positively
or negatively charged with  respect to the solution. A potential difference
develops between the electrode and the electrolyte which is called
electrode potential. When the concentrations of all the species involved
in a half-cell is unity then the electrode potential is known as standard
electrode potential. According to IUPAC convention, standard
reduction potentials are now called standard electrode potentials. In a
galvanic cell, the half-cell in which oxidation takes place is called anode
and it has a negative potential with respect to the solution. The other
half-cell in which reduction takes place is called cathode and it has a
positive potential with respect to the solution. Thus, there exists a
potential difference between the two electrodes and as soon as the
switch is in the on position the electrons flow from negative electrode
to positive electrode. The direction of current flow is opposite to that of
electron flow.
2.2 Galvanic Cells 2.2 Galvanic Cells 2.2 Galvanic Cells 2.2 Galvanic Cells 2.2 Galvanic Cells
Reprint 2024-25
Page 4


Electrochemistry is the study of production of
electricity from energy released during spontaneous
chemical reactions  and the use of electrical energy
to bring about non-spontaneous chemical
transformations. The subject is of importance both
for theoretical and practical considerations. A large
number of metals, sodium hydroxide, chlorine,
fluorine and many other chemicals are produced by
electrochemical methods. Batteries and fuel cells
convert chemical energy into electrical energy and are
used on a large scale in various instruments and
devices. The reactions carried out electrochemically
can be energy efficient and less polluting. Therefore,
study of electrochemistry is important for creating new
technologies that are ecofriendly. The transmission of
sensory signals through cells to brain and vice versa
and communication between the cells are known to
have electrochemical origin. Electrochemistry, is
therefore, a very vast and interdisciplinary subject. In
this Unit, we will cover only some of its important
elementary aspects.
After studying this Unit, you will be
able to
· describe an electrochemical cell
and differentiate between galvanic
and electrolytic cells;
· apply Nernst equation for
calculating the emf of galvanic cell
and define standard potential of
the cell;
· derive relation between standard
potential of the cell, Gibbs energy
of cell reaction and its equilibrium
constant;
· define resistivity (r), conductivity
(k) and molar conductivity (?
m
) of
ionic solutions;
· differentiate between ionic
(electrolytic) and electronic
conductivity;
· describe the method for
measurement of conductivity of
electrolytic solutions and
calculation of their molar
conductivity;
· justify the variation of
conductivity and molar
conductivity of solutions with
change in their concentration and
define ?
m
? (molar conductivity at
zero concentration or infinite
dilution);
· enunciate Kohlrausch law and
learn its applications;
· understand quantitative aspects
of electrolysis;
· describe the construction of some
primary and secondary batteries
and fuel cells;
· explain corrosion as an
electrochemical process.
Objectives
Chemical reactions can be used to produce electrical energy,
conversely, electrical energy can be used to carry out chemical
reactions that do not proceed spontaneously.
2
Electrochemistry
Unit Unit Unit Unit Unit
2
Electrochemistry
Reprint 2024-25
32 Chemistry
Cu
E
ext
>1.1
e
–
Current
Cathode
+ve
Anode
–ve
Zn
Fig. 2.2
Functioning of Daniell
cell when external
voltage E
ext
 opposing the
cell potential is applied.
We had studied the construction and functioning of Daniell cell
(Fig. 2.1). This cell converts the chemical energy liberated during the
redox reaction
Zn(s) + Cu
2+
(aq) ® Zn
2+
(aq) + Cu(s) (2.1)
to electrical energy and has an electrical
potential equal to 1.1 V when concentration
of Zn
2+ 
and Cu
2+
 ions is unity (1 mol dm
–3
)
*
.
Such a device is called a galvanic or a
voltaic cell.
If an external opposite potential is applied
in the galvanic cell [Fig. 2.2(a)] and increased
slowly, we find that the reaction continues to
take place till the opposing voltage reaches
the value 1.1 V [Fig. 2.2(b)] when, the reaction
stops altogether and no current flows through
the cell. Any further increase in the external
potential again starts the reaction but in the
opposite direction [Fig. 2.2(c)]. It now functions
as an electrolytic cell, a device for using
electrical energy to carry non-spontaneous
chemical reactions. Both types of cells are
quite important and we shall study some of
their salient features in the following pages.
*Strictly speaking activity should be used instead of concentration. It is directly proportional to concentration. In dilute
solutions, it is equal to concentration. You will study more about it in higher classes.
2.1 2.1 2.1 2.1 2.1 Electrochemical Electrochemical Electrochemical Electrochemical Electrochemical
Cells Cells Cells Cells Cells
Fig. 2.1: Daniell cell having electrodes of zinc and
copper dipping in the solutions of their
respective salts.
salt
bridge
Zn Cu
anode
cathode current
ZnSO
4
CuSO
4
E <
ext
1.1V
e
-ve +ve
I=0
Zn Cu
ZnSO
4
CuSO
4
E =
ext
1.1V
When E
ext
 < 1.1 V
(i) Electrons flow from Zn rod to
Cu rod hence current flows
from Cu to Zn.
(ii) Zn dissolves at anode and
copper deposits at cathode.
When E
ext
 = 1.1 V
(i) No flow of
electrons or
current.
(ii) No chemical
reaction.
When E
ext
 > 1.1 V
(i) Electrons flow
from Cu to Zn
and current flows
from Zn to Cu.
(ii) Zinc is deposited
at the zinc
electrode and
copper dissolves at
copper electrode.
(a)
(b)
(c)
Reprint 2024-25
33 Electrochemistry
As mentioned earlier a galvanic cell is an electrochemical cell that
converts the chemical energy of a spontaneous redox reaction into
electrical energy. In this device the Gibbs energy of the spontaneous
redox reaction is converted into electrical work which may be used for
running a motor or other electrical gadgets  like heater, fan,
geyser, etc.
Daniell cell discussed earlier is one such cell in which the following
redox reaction occurs.
Zn(s) + Cu
2+
(aq) ® Zn
2+
 (aq) + Cu(s)
This reaction is a combination of two half reactions whose addition
gives the overall cell reaction:
(i) Cu
2+
   +  2e
–
    ® Cu(s) (reduction half reaction) (2.2)
(ii) Zn(s)  ® Zn
2+
 + 2e
–
(oxidation half reaction) (2.3)
These reactions occur in two different portions of the Daniell cell.
The reduction half reaction occurs on the copper electrode while the
oxidation half reaction occurs on the zinc electrode. These two portions
of the cell are also called half-cells or redox couples. The copper
electrode may be called the reduction half cell and the zinc electrode,
the oxidation half-cell.
We can construct innumerable number of galvanic cells on the pattern
of Daniell cell by taking combinations of different half-cells. Each half-
cell consists of a metallic electrode dipped into an electrolyte. The two
half-cells are connected by a metallic wire through a voltmeter and a
switch externally. The electrolytes of the two half-cells are connected
internally through a salt bridge as shown in Fig. 2.1.  Sometimes, both
the electrodes dip in the same electrolyte solution and in such cases we
do not require a salt bridge.
At each electrode-electrolyte interface there is a tendency of metal
ions from the solution to deposit on the metal electrode trying to make
it positively charged. At the same time, metal atoms of the electrode
have a tendency to go into the solution as ions and leave behind the
electrons at the electrode trying to make it negatively charged. At
equilibrium, there is a separation of charges and depending on the
tendencies of the two opposing reactions, the electrode may be positively
or negatively charged with  respect to the solution. A potential difference
develops between the electrode and the electrolyte which is called
electrode potential. When the concentrations of all the species involved
in a half-cell is unity then the electrode potential is known as standard
electrode potential. According to IUPAC convention, standard
reduction potentials are now called standard electrode potentials. In a
galvanic cell, the half-cell in which oxidation takes place is called anode
and it has a negative potential with respect to the solution. The other
half-cell in which reduction takes place is called cathode and it has a
positive potential with respect to the solution. Thus, there exists a
potential difference between the two electrodes and as soon as the
switch is in the on position the electrons flow from negative electrode
to positive electrode. The direction of current flow is opposite to that of
electron flow.
2.2 Galvanic Cells 2.2 Galvanic Cells 2.2 Galvanic Cells 2.2 Galvanic Cells 2.2 Galvanic Cells
Reprint 2024-25
34 Chemistry
The potential difference between the two electrodes of a galvanic
cell is called the cell potential and is measured in volts. The cell
potential is the difference between the electrode potentials (reduction
potentials) of the cathode and anode. It is called the cell electromotive
force (emf) of the cell when no current is drawn through the cell. It
is now an accepted convention that we keep the anode on the left and
the cathode on the right  while representing the galvanic cell. A galvanic
cell is generally represented by putting a vertical line between metal
and electrolyte solution and putting a double vertical line between
the two electrolytes connected by a salt bridge. Under this convention
the emf of the cell is positive and is given by the potential of the half-
cell on the right hand side minus the potential of the half-cell on the
left hand side i.e.,
E
cell 
= E
right  
– E
left
This is illustrated by the following example:
Cell reaction:
Cu(s) + 2Ag
+
(aq) ¾® Cu
2+
(aq) + 2 Ag(s) (2.4)
Half-cell reactions:
Cathode (reduction):   2Ag
+
(aq)
 
+ 2e
–
 ® 2Ag(s) (2.5)
Anode (oxidation):    Cu(s)  ® Cu
2+
(aq) + 2e
–
(2.6)
It can be seen that the sum of (3.5) and (3.6) leads to overall reaction
(2.4) in the cell and that silver electrode acts as a cathode and copper
electrode acts as an anode. The cell can be represented as:
Cu(s)|Cu
2+
(aq)| |Ag
+
(aq)|Ag(s)
and we have E
cell
 = E
right
 – E
left
 = E
Ag
+
úAg
 – E
Cu
2+
úCu
(2.7)
The potential of individual half-cell cannot be measured. We can
measure only the difference between the two half-cell potentials that
gives the emf of the cell. If we arbitrarily choose the potential of one
electrode (half-cell) then that of the other can be determined with respect
to this. According to convention, a half-cell
called standard hydrogen electrode (Fig.3.3)
represented by Pt(s)ú H
2
(g)ú H
+
(aq), is assigned
a zero potential at all temperatures
corresponding to the reaction
    H
+
 (aq) + e
–
  ®  
1
2
H
2
(g)
The standard hydrogen electrode consists
of a platinum electrode coated with platinum
black. The electrode is dipped in an acidic
solution and pure hydrogen gas is bubbled
through it.  The concentration of both the
reduced and oxidised forms of hydrogen is
maintained at unity (Fig. 2.3). This implies
that the pressure of hydrogen gas is one bar
and the concentration of hydrogen ion in the
solution is one molar.
2.2.1
Measurement
of Electrode
Potential
Fig. 2.3: Standard Hydrogen Electrode (SHE).
Reprint 2024-25
Page 5


Electrochemistry is the study of production of
electricity from energy released during spontaneous
chemical reactions  and the use of electrical energy
to bring about non-spontaneous chemical
transformations. The subject is of importance both
for theoretical and practical considerations. A large
number of metals, sodium hydroxide, chlorine,
fluorine and many other chemicals are produced by
electrochemical methods. Batteries and fuel cells
convert chemical energy into electrical energy and are
used on a large scale in various instruments and
devices. The reactions carried out electrochemically
can be energy efficient and less polluting. Therefore,
study of electrochemistry is important for creating new
technologies that are ecofriendly. The transmission of
sensory signals through cells to brain and vice versa
and communication between the cells are known to
have electrochemical origin. Electrochemistry, is
therefore, a very vast and interdisciplinary subject. In
this Unit, we will cover only some of its important
elementary aspects.
After studying this Unit, you will be
able to
· describe an electrochemical cell
and differentiate between galvanic
and electrolytic cells;
· apply Nernst equation for
calculating the emf of galvanic cell
and define standard potential of
the cell;
· derive relation between standard
potential of the cell, Gibbs energy
of cell reaction and its equilibrium
constant;
· define resistivity (r), conductivity
(k) and molar conductivity (?
m
) of
ionic solutions;
· differentiate between ionic
(electrolytic) and electronic
conductivity;
· describe the method for
measurement of conductivity of
electrolytic solutions and
calculation of their molar
conductivity;
· justify the variation of
conductivity and molar
conductivity of solutions with
change in their concentration and
define ?
m
? (molar conductivity at
zero concentration or infinite
dilution);
· enunciate Kohlrausch law and
learn its applications;
· understand quantitative aspects
of electrolysis;
· describe the construction of some
primary and secondary batteries
and fuel cells;
· explain corrosion as an
electrochemical process.
Objectives
Chemical reactions can be used to produce electrical energy,
conversely, electrical energy can be used to carry out chemical
reactions that do not proceed spontaneously.
2
Electrochemistry
Unit Unit Unit Unit Unit
2
Electrochemistry
Reprint 2024-25
32 Chemistry
Cu
E
ext
>1.1
e
–
Current
Cathode
+ve
Anode
–ve
Zn
Fig. 2.2
Functioning of Daniell
cell when external
voltage E
ext
 opposing the
cell potential is applied.
We had studied the construction and functioning of Daniell cell
(Fig. 2.1). This cell converts the chemical energy liberated during the
redox reaction
Zn(s) + Cu
2+
(aq) ® Zn
2+
(aq) + Cu(s) (2.1)
to electrical energy and has an electrical
potential equal to 1.1 V when concentration
of Zn
2+ 
and Cu
2+
 ions is unity (1 mol dm
–3
)
*
.
Such a device is called a galvanic or a
voltaic cell.
If an external opposite potential is applied
in the galvanic cell [Fig. 2.2(a)] and increased
slowly, we find that the reaction continues to
take place till the opposing voltage reaches
the value 1.1 V [Fig. 2.2(b)] when, the reaction
stops altogether and no current flows through
the cell. Any further increase in the external
potential again starts the reaction but in the
opposite direction [Fig. 2.2(c)]. It now functions
as an electrolytic cell, a device for using
electrical energy to carry non-spontaneous
chemical reactions. Both types of cells are
quite important and we shall study some of
their salient features in the following pages.
*Strictly speaking activity should be used instead of concentration. It is directly proportional to concentration. In dilute
solutions, it is equal to concentration. You will study more about it in higher classes.
2.1 2.1 2.1 2.1 2.1 Electrochemical Electrochemical Electrochemical Electrochemical Electrochemical
Cells Cells Cells Cells Cells
Fig. 2.1: Daniell cell having electrodes of zinc and
copper dipping in the solutions of their
respective salts.
salt
bridge
Zn Cu
anode
cathode current
ZnSO
4
CuSO
4
E <
ext
1.1V
e
-ve +ve
I=0
Zn Cu
ZnSO
4
CuSO
4
E =
ext
1.1V
When E
ext
 < 1.1 V
(i) Electrons flow from Zn rod to
Cu rod hence current flows
from Cu to Zn.
(ii) Zn dissolves at anode and
copper deposits at cathode.
When E
ext
 = 1.1 V
(i) No flow of
electrons or
current.
(ii) No chemical
reaction.
When E
ext
 > 1.1 V
(i) Electrons flow
from Cu to Zn
and current flows
from Zn to Cu.
(ii) Zinc is deposited
at the zinc
electrode and
copper dissolves at
copper electrode.
(a)
(b)
(c)
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33 Electrochemistry
As mentioned earlier a galvanic cell is an electrochemical cell that
converts the chemical energy of a spontaneous redox reaction into
electrical energy. In this device the Gibbs energy of the spontaneous
redox reaction is converted into electrical work which may be used for
running a motor or other electrical gadgets  like heater, fan,
geyser, etc.
Daniell cell discussed earlier is one such cell in which the following
redox reaction occurs.
Zn(s) + Cu
2+
(aq) ® Zn
2+
 (aq) + Cu(s)
This reaction is a combination of two half reactions whose addition
gives the overall cell reaction:
(i) Cu
2+
   +  2e
–
    ® Cu(s) (reduction half reaction) (2.2)
(ii) Zn(s)  ® Zn
2+
 + 2e
–
(oxidation half reaction) (2.3)
These reactions occur in two different portions of the Daniell cell.
The reduction half reaction occurs on the copper electrode while the
oxidation half reaction occurs on the zinc electrode. These two portions
of the cell are also called half-cells or redox couples. The copper
electrode may be called the reduction half cell and the zinc electrode,
the oxidation half-cell.
We can construct innumerable number of galvanic cells on the pattern
of Daniell cell by taking combinations of different half-cells. Each half-
cell consists of a metallic electrode dipped into an electrolyte. The two
half-cells are connected by a metallic wire through a voltmeter and a
switch externally. The electrolytes of the two half-cells are connected
internally through a salt bridge as shown in Fig. 2.1.  Sometimes, both
the electrodes dip in the same electrolyte solution and in such cases we
do not require a salt bridge.
At each electrode-electrolyte interface there is a tendency of metal
ions from the solution to deposit on the metal electrode trying to make
it positively charged. At the same time, metal atoms of the electrode
have a tendency to go into the solution as ions and leave behind the
electrons at the electrode trying to make it negatively charged. At
equilibrium, there is a separation of charges and depending on the
tendencies of the two opposing reactions, the electrode may be positively
or negatively charged with  respect to the solution. A potential difference
develops between the electrode and the electrolyte which is called
electrode potential. When the concentrations of all the species involved
in a half-cell is unity then the electrode potential is known as standard
electrode potential. According to IUPAC convention, standard
reduction potentials are now called standard electrode potentials. In a
galvanic cell, the half-cell in which oxidation takes place is called anode
and it has a negative potential with respect to the solution. The other
half-cell in which reduction takes place is called cathode and it has a
positive potential with respect to the solution. Thus, there exists a
potential difference between the two electrodes and as soon as the
switch is in the on position the electrons flow from negative electrode
to positive electrode. The direction of current flow is opposite to that of
electron flow.
2.2 Galvanic Cells 2.2 Galvanic Cells 2.2 Galvanic Cells 2.2 Galvanic Cells 2.2 Galvanic Cells
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34 Chemistry
The potential difference between the two electrodes of a galvanic
cell is called the cell potential and is measured in volts. The cell
potential is the difference between the electrode potentials (reduction
potentials) of the cathode and anode. It is called the cell electromotive
force (emf) of the cell when no current is drawn through the cell. It
is now an accepted convention that we keep the anode on the left and
the cathode on the right  while representing the galvanic cell. A galvanic
cell is generally represented by putting a vertical line between metal
and electrolyte solution and putting a double vertical line between
the two electrolytes connected by a salt bridge. Under this convention
the emf of the cell is positive and is given by the potential of the half-
cell on the right hand side minus the potential of the half-cell on the
left hand side i.e.,
E
cell 
= E
right  
– E
left
This is illustrated by the following example:
Cell reaction:
Cu(s) + 2Ag
+
(aq) ¾® Cu
2+
(aq) + 2 Ag(s) (2.4)
Half-cell reactions:
Cathode (reduction):   2Ag
+
(aq)
 
+ 2e
–
 ® 2Ag(s) (2.5)
Anode (oxidation):    Cu(s)  ® Cu
2+
(aq) + 2e
–
(2.6)
It can be seen that the sum of (3.5) and (3.6) leads to overall reaction
(2.4) in the cell and that silver electrode acts as a cathode and copper
electrode acts as an anode. The cell can be represented as:
Cu(s)|Cu
2+
(aq)| |Ag
+
(aq)|Ag(s)
and we have E
cell
 = E
right
 – E
left
 = E
Ag
+
úAg
 – E
Cu
2+
úCu
(2.7)
The potential of individual half-cell cannot be measured. We can
measure only the difference between the two half-cell potentials that
gives the emf of the cell. If we arbitrarily choose the potential of one
electrode (half-cell) then that of the other can be determined with respect
to this. According to convention, a half-cell
called standard hydrogen electrode (Fig.3.3)
represented by Pt(s)ú H
2
(g)ú H
+
(aq), is assigned
a zero potential at all temperatures
corresponding to the reaction
    H
+
 (aq) + e
–
  ®  
1
2
H
2
(g)
The standard hydrogen electrode consists
of a platinum electrode coated with platinum
black. The electrode is dipped in an acidic
solution and pure hydrogen gas is bubbled
through it.  The concentration of both the
reduced and oxidised forms of hydrogen is
maintained at unity (Fig. 2.3). This implies
that the pressure of hydrogen gas is one bar
and the concentration of hydrogen ion in the
solution is one molar.
2.2.1
Measurement
of Electrode
Potential
Fig. 2.3: Standard Hydrogen Electrode (SHE).
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35 Electrochemistry
At 298 K the emf of the cell, standard hydrogen electrode ççsecond
half-cell constructed by taking standard hydrogen electrode as anode
(reference half-cell) and the other half-cell  as cathode, gives the reduction
potential of the other half-cell. If the concentrations of the oxidised and
the reduced forms of the species in the right hand half-cell are unity,
then the cell potential is equal to standard electrode potential, E
o
R
 of
the given half-cell.
E
o
 = E
o
R
 – E
o
L
As E
o
L
 for standard hydrogen electrode is zero.
E
o
 = E
o
R
 – 0 = E
o
R
The measured emf of the cell:
Pt(s) ç H
2
(g, 1 bar) ç H
+ 
(aq, 1 M) çç Cu
2+
 (aq, 1 M) ú Cu
is 0.34 V and it is also the value for the standard electrode potential
of the half-cell corresponding to the reaction:
Cu
2+
 (aq, 1M) + 2 e
–
  ®  Cu(s)
Similarly, the measured emf of the cell:
Pt(s) ç H
2
(g, 1 bar) ç H
+ 
(aq, 1 M) çç Zn
2+
 (aq, 1M) ç Zn
is -0.76 V corresponding to the standard electrode potential of the
half-cell reaction:
Zn
2+
 (aq, 1 M) + 2e
–
 ®  Zn(s)
The positive  value of the standard electrode potential in the first
case indicates that Cu
2+
 ions get reduced more easily than H
+
 ions. The
reverse process cannot occur, that is, hydrogen ions cannot oxidise Cu
(or alternatively we can say that hydrogen gas can reduce copper ion)
under the standard conditions described above. Thus, Cu does not
dissolve in HCl.  In nitric acid it is oxidised by nitrate ion and not by
hydrogen ion.  The negative value of the standard electrode potential
in the second case indicates that hydrogen ions can oxidise zinc (or
zinc can reduce hydrogen ions).
In view of this convention, the half reaction for the Daniell cell in
Fig. 2.1 can be written as:
Left electrode: Zn(s) ® Zn
2+
 (aq, 1 M) + 2 e
–
Right electrode: Cu
2+
 (aq, 1 M) + 2 e
–
 ® Cu(s)
The overall reaction of the cell is the sum of above two reactions
and we obtain the equation:
Zn(s) + Cu
2+
 (aq) ® Zn
2+ 
(aq) + Cu(s)
emf of the cell = E
o
cell
= E
o
R
 – E
o
L
= 0.34V – (– 0.76)V = 1.10 V
Sometimes metals like platinum or gold are used as inert electrodes.
They do not participate in the reaction but provide their surface for
oxidation or reduction reactions and for the conduction of electrons.
For example, Pt is used in the following half-cells:
Hydrogen electrode:  Pt(s)|H
2
(g)| H
+
(aq)
With half-cell reaction:  H
+ 
(aq)+ e
– 
® ½ H
2
(g)
Bromine electrode: Pt(s)|Br
2
(aq)| Br
–
(aq)
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