Electrochemistry (NCERT) Notes | EduRev

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: Electrochemistry (NCERT) Notes | EduRev

 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 (?), conductivity
(?) 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.
3
Electr Electr Electr Electr Electrochemistr ochemistr ochemistr ochemistr ochemistry y y y y
Unit Unit Unit Unit Unit
3
Electr Electr Electr Electr Electrochemistr ochemistr ochemistr ochemistr ochemistry y y y y
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 (?), conductivity
(?) 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.
3
Electr Electr Electr Electr Electrochemistr ochemistr ochemistr ochemistr ochemistry y y y y
Unit Unit Unit Unit Unit
3
Electr Electr Electr Electr Electrochemistr ochemistr ochemistr ochemistr ochemistry y y y y
64 Chemistry
In Class XI, Unit 8, we had studied the construction and functioning
of Daniell cell (Fig. 3.1). This cell converts the chemical energy liberated
during the redox reaction
Zn(s) + Cu
2+
(aq) ? Zn
2+
(aq) + Cu(s) (3.1)
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 [Fig. 3.2(a)] and increased slowly,
we find that the reaction continues to take
place till the opposing voltage reaches the
value 1.1 V [Fig. 3.2(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.
3.2(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.
3.1 Electrochemical
Cells
Fig. 3.1: Daniell cell having electrodes of zinc and
copper dipping in the solutions of their
respective salts.
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 (?), conductivity
(?) 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.
3
Electr Electr Electr Electr Electrochemistr ochemistr ochemistr ochemistr ochemistry y y y y
Unit Unit Unit Unit Unit
3
Electr Electr Electr Electr Electrochemistr ochemistr ochemistr ochemistr ochemistry y y y y
64 Chemistry
In Class XI, Unit 8, we had studied the construction and functioning
of Daniell cell (Fig. 3.1). This cell converts the chemical energy liberated
during the redox reaction
Zn(s) + Cu
2+
(aq) ? Zn
2+
(aq) + Cu(s) (3.1)
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 [Fig. 3.2(a)] and increased slowly,
we find that the reaction continues to take
place till the opposing voltage reaches the
value 1.1 V [Fig. 3.2(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.
3.2(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.
3.1 Electrochemical
Cells
Fig. 3.1: Daniell cell having electrodes of zinc and
copper dipping in the solutions of their
respective salts.
65 Electrochemistry
As mentioned earlier (Class XI, Unit 8) 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) (3.2)
(ii) Zn(s)  ? Zn
2+
 + 2e
–
(oxidation half reaction) (3.3)
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. 3.1.  Sometimes, both
the electrodes dip in the same electrolyte solution and in such cases we
don’t require a salt bridge.
Fig. 3.2: Functioning of Daniell cell when external voltage E
ext
 opposing the
cell potential is applied.
3.2 Galvanic Cells
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 (?), conductivity
(?) 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.
3
Electr Electr Electr Electr Electrochemistr ochemistr ochemistr ochemistr ochemistry y y y y
Unit Unit Unit Unit Unit
3
Electr Electr Electr Electr Electrochemistr ochemistr ochemistr ochemistr ochemistry y y y y
64 Chemistry
In Class XI, Unit 8, we had studied the construction and functioning
of Daniell cell (Fig. 3.1). This cell converts the chemical energy liberated
during the redox reaction
Zn(s) + Cu
2+
(aq) ? Zn
2+
(aq) + Cu(s) (3.1)
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 [Fig. 3.2(a)] and increased slowly,
we find that the reaction continues to take
place till the opposing voltage reaches the
value 1.1 V [Fig. 3.2(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.
3.2(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.
3.1 Electrochemical
Cells
Fig. 3.1: Daniell cell having electrodes of zinc and
copper dipping in the solutions of their
respective salts.
65 Electrochemistry
As mentioned earlier (Class XI, Unit 8) 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) (3.2)
(ii) Zn(s)  ? Zn
2+
 + 2e
–
(oxidation half reaction) (3.3)
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. 3.1.  Sometimes, both
the electrodes dip in the same electrolyte solution and in such cases we
don’t require a salt bridge.
Fig. 3.2: Functioning of Daniell cell when external voltage E
ext
 opposing the
cell potential is applied.
3.2 Galvanic Cells
66 Chemistry
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.
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) ?? Cu
2+
(aq) + 2 Ag(s) (3.4)
Half-cell reactions:
Cathode (reduction):   2Ag
+
(aq)
 
+ 2e
–
 ? 2Ag(s) (3.5)
Anode (oxidation):    Cu(s)  ? Cu
2+
(aq) + 2e
–
(3.6)
It can be seen that the sum of (3.5) and (3.6) leads to overall reaction
(3.4) 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
+
? Ag
 – E
Cu
2+
? Cu
(3.7)
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-
3.2.1 Measurement
of Electrode
Potential
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 (?), conductivity
(?) 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.
3
Electr Electr Electr Electr Electrochemistr ochemistr ochemistr ochemistr ochemistry y y y y
Unit Unit Unit Unit Unit
3
Electr Electr Electr Electr Electrochemistr ochemistr ochemistr ochemistr ochemistry y y y y
64 Chemistry
In Class XI, Unit 8, we had studied the construction and functioning
of Daniell cell (Fig. 3.1). This cell converts the chemical energy liberated
during the redox reaction
Zn(s) + Cu
2+
(aq) ? Zn
2+
(aq) + Cu(s) (3.1)
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 [Fig. 3.2(a)] and increased slowly,
we find that the reaction continues to take
place till the opposing voltage reaches the
value 1.1 V [Fig. 3.2(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.
3.2(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.
3.1 Electrochemical
Cells
Fig. 3.1: Daniell cell having electrodes of zinc and
copper dipping in the solutions of their
respective salts.
65 Electrochemistry
As mentioned earlier (Class XI, Unit 8) 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) (3.2)
(ii) Zn(s)  ? Zn
2+
 + 2e
–
(oxidation half reaction) (3.3)
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. 3.1.  Sometimes, both
the electrodes dip in the same electrolyte solution and in such cases we
don’t require a salt bridge.
Fig. 3.2: Functioning of Daniell cell when external voltage E
ext
 opposing the
cell potential is applied.
3.2 Galvanic Cells
66 Chemistry
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.
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) ?? Cu
2+
(aq) + 2 Ag(s) (3.4)
Half-cell reactions:
Cathode (reduction):   2Ag
+
(aq)
 
+ 2e
–
 ? 2Ag(s) (3.5)
Anode (oxidation):    Cu(s)  ? Cu
2+
(aq) + 2e
–
(3.6)
It can be seen that the sum of (3.5) and (3.6) leads to overall reaction
(3.4) 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
+
? Ag
 – E
Cu
2+
? Cu
(3.7)
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-
3.2.1 Measurement
of Electrode
Potential
67 Electrochemistry
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)? H
2
(g)? 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. 3.3). This implies
that the pressure of hydrogen gas is one
bar and the concentration of hydrogen ion
in the solution is one molar.
At 298 K the emf of the cell, standard
hydrogen electrode ?? second 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

R
 of the given half-cell.
E

 = E

R
  - E

L
As E

L
 for standard hydrogen electrode
is zero.
E

 =  E

R
  – 0  =  E

R
The measured emf of the cell :
Pt(s) ? H
2
(g, 1 bar) ? H
+ 
(aq, 1 M) ?? Cu
2+
 (aq, 1 M)? 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) ? H
2
(g, 1 bar) ? H
+ 
(aq, 1 M) ?? Zn
2+
 (aq, 1M) ? 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).
Fig. 3.3: Standard Hydrogen Electrode (SHE).
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