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Electrochemistry is the branch of physical
chemistry which deals with the relationship between
electrical energy and chemical changes taking place in
redox reactions
Electrolytes and Electrolysis
(1) Definition : “ The substances whose aqueous
solution undergo decomposition into ions when electric
current is passed through them are known as
electrolytes and the whole process is known as
electrolysis or electrolytic decomposition .”
Solutions of acids, bases, salts in water and fused
salts etc. are the examples of electrolytes. Electrolytes
may be weak or strong. Solutions of cane sugar,
glycerine, alcohol etc., are examples of non-
electrolytes.
(2) Electrolytic cell or Voltameter : The device
in which the process of electrolysis or electrolytic
decomposition is carried out is known as electrolytic
cell or voltameter.
(i) Voltameter convert electrical energy into
chemical energy.
(ii) The electrode on which oxidation takes place
is called anode (or +ve pole) and the electrode on
which reduction takes place is called cathode (or – ve
pole)
(iii) During electrolysis in voltameter cations are
discharged on cathode and anions on anode.
(iv) In voltameter, outside the electrolyte
electrons flow from anode to cathode and current flow
from cathode to anode.
For voltameter, Ecell veandΔG ve.
(v) The anions on reaching the anode give up
their electrons and converted into the neutral atoms.
At anode : A A e
(Oxidation)
(vi) On the other hand cations on reaching the
cathode take up electrons supplied by battery and
converted to the neutral atoms.
At cathode :B B (^) e (Reduction)
This overall change is known as primary change
and products formed is known as primary products.
The primary products may be collected as such or
they undergo further change to form molecules or
compounds. These are called secondary products and
the change is known as secondary change.
(3) Preferential discharge theory : According to
this theory “ If more than one type of ion is attracted
towards a particular electrode, then the ion is discharged
one which requires least energy or ions with lower
discharge potential or which occur low in the
electrochemical series ”.
The potential at which the ion is discharged or
deposited on the appropriate electrode is termed the
Flow of electrons
Flow of current
Anode Cathod
e
Electrochemistry
Chapter
discharge or deposition potential , (D.P.). The values of
discharge potential are different for different ions.
The decreasing order of discharge potential or the
increasing order of deposition of some of the ions is
given below,
For cations : , , , , , , ,
2 2 3 2 Li K Na Ca Mg Al Zn
2 Fe , , , , ,.
2 2 2 3 Ni H Cu Hg Ag Au
For anions : , 3 , , , ,.
2 4
SO NO OH Cl Br I
Table : 12.1 Products of electrolysis of some electrolytes
Electrolyte Electrode Product at cathode Product at anode
Aqueous NaOH Pt or Graphite
2 H 2 e H 2 ^ OH O HO 2 e 2
1 (^2 )
Fused NaOH Pt or Graphite Na e Na OH O HO 2 e 2
1 (^2 )
Aqueous NaCl Pt or Graphite 2 H 2 e H 2
2 Cl Cl 2 2 e
Fused NaCl Pt or Graphite Na e Na 2 Cl ^ Cl 2 e 2
Aqueous CuSO 4 Pt or Graphite Cu^2 ^ 2 e Cu OH O HO 2 e 2
1 2 2 2
Aqueous CuSO 4 Cu electrode (^) Cu^2 ^ 2 e Cu Cu oxidised to Cu^2 ions
Dilute H 2 SO 4 Pt electrode 2 H 2 e H 2
OH O HO 2 e 2
1 (^2 )
Conc. H 2 SO 4 Pt electrode 2 H 2 e H 2 Peroxodisulphuric
acid( H 2 S 2 O 8 )
Aqueous AgNO 3 Pt electrode Ag e Ag OH O HO 2 e 2
1 (^2 )
Aqueous AgNO 3 Ag electrode (^) Ag e Ag Ag oxidised to Ag ions
(4) Application of electrolysis : Electrolysis has
wide applications in industries. Some of the important
applications are, as follows,
(i) Production of hydrogen by electrolysis of
water.
(ii) Manufacture of heavy water ( D 2 O ).
(iii) The metals like Na , K , Mg , Al , etc., are
obtained by electrolysis of fused electrolytes.
(iv) Non-metals like hydrogen, fluorine, chlorine
are obtained by electrolysis.
(v) In this method pure metal is deposited at
cathode from a solution containing the metal ions
Ag , Cu etc.
(vi) Compounds like NaOH, KOH,
Na 2 CO 3 , KClO 3 ,^ white^ lead,^ KMnO 4 etc. are synthesised
by electrosynthesis method.
(vii) Electroplating : The process of coating an
inferior metal with a superior metal by electrolysis is
known as electroplating. The aim of electroplating is, to
prevent the inferior metal from corrosion and to make
it more attractive in appearance. The object to be
plated is made the cathode of an electrolytic cell that
contains a solution of ions of the metal to be deposited.
For Anode Cathode Electrolyte
electroplati
ng
With copper Cu Object CuSO (^) 4 dilute H 2 SO 4
With silver Ag Object K [ Ag ( CN ) 2 ]
With nickel Ni Object Nickel
ammonium
sulphate
With gold Au Object K [ Au ( CN ) 2 ]
With zinc Zn Iron
objects
ZnSO 4
With tin Sn Iron
objects
SnSO 4
Thickness of coated layer : Let the dimensions of
metal sheet to be coated be( a cm bcm ).
Thickness of coated layer ccm
Volume of coated layer
3 ( a b c ) cm
Mass of the deposited substanceVolume density
( a b c ) dg
96500
( )
I t E a b c d
Using above relation we may calculate the
thickness of coated layer.
transfer of matter. matter in the form of ions.
(iv) Conductivity decreases
with increase in temperature.
(iv) Conductivity increases
with increases in temperature and degree of
hydration due to decreases in viscosity of medium.
The electrolyte may, therefore, be defined as the
substance whose aqueous solution or fused state
conduct electricity accompanied by chemical
decomposition. The conduction of current through
electrolyte is due to the movement of ions.
On the contrary, substances, which in the form of
their solutions or in their molten state do not conduct
electricity, are called non-electrolytes.
Electrolytic conduction
When a voltage is applied to the electrodes dipped
into an electrolytic solution, ions of the electrolyte
move and, therefore, electric current flows through the
electrolytic solution. The power of the electrolytes to
conduct electric current is termed conductance or
conductivity.
(1) Ohm's law : This law states that the current
flowing through a conductor is directly proportional to
the potential difference across it, i.e.,I V
where I is the current strength (In Amperes) and V
is the potential difference applied across the conductor
(In Volts)
or R
V I or V IR
where R is the constant of proportionality and is
known as resistance of the conductor. It is expressed in
Ohm's and is represented as .The above equation is
known as Ohm's law. Ohm's law may also be stated as,
“ the strength of current flowing through a
conductor is directly proportional to the potential
difference applied across the conductor and inversely
proportional to the resistance of the conductor .”
(2) Resistance : It measures the obstruction to
the flow of current. The resistance of any conductor is
directly proportional to the length ( l ) and inversely
proportional to the area of cross-section ( a ) so that
a
l R ρ a
l R or
where (rho) is the constant of proportionality
and is called specific resistance or resistivity. The
resistance depends upon the nature of the material.
Units : The unit of resistance is ohm ().In terms
of SI, base unit is equal to( )/( ).
2 3 2 kgm s A
(3) Resistivity or specific resistance : We know
that resistance R is
a
l R ; Now, if 2 l 1 cm , a 1 cm then R
Thus, resistivity is defined as the resistance of a
conductor of 1 cm length and having area of cross-
section equal to 1.
2 cm
Units : The units of resistivity are
cm
cm Ohm l
a R
2 . Ohm. cm
Its SI units are Ohm metre ( m ).But quite often
Ohm centimetre ( cm )is also used.
(4) Conductance : It is a measure of the ease with
which current flows through a conductor. It is an
additive property. It is expressed as G. It is reciprocal
of the resistance, i.e.,
R
G
1
Units : The units of conductance are reciprocal
Ohm( )
1 ohm or mho. Ohm is also abbreviated as so
that
1 Ohm may be written as.
1
According to SI system, the units of electrical
conductance is Siemens, S ( i.e., 1 S 1 ).
1
(5) Conductivity : The inverse of resistivity is
called conductivity (or specific conductance). It is
represented by the symbol, (Greek kappa). The
IUPAC has recommended the use of term conductivity
over specific conductance. It may be defined as, the
conductance of a solution of 1 cm length and having 1 sq.
cm as the area of cross-section. In other words,
conductivity is the conductance of one centimetre cube of
a solution of an electrolyte.
Thus,
Units : The units of conductivity are
1
.
Ohm Ohmcm
cm
1 1 cm
In SI units, l is expressed in m area of cross-
section in
2 m so that the units of conductivity are
S.
1 m
(6) Molar conductivity or molar conductance :
Molar conductivity is defined as the conducting power
of all the ions produced by dissolving one mole of an
electrolyte in solution.
It is denoted by (lambda). Molar conductance is
related to specific conductance ( ) as,
M
where, M is the molar concentration.
If M is in the units of molarity i.e., moles per litre
( ),
1 mol L the may be expressed as,
M
1000
For the solution containing 1 gm mole of
electrolyte placed between two parallel electrodes of 1
sq. cm area of cross-section and one cm apart,
Conductanc e( G )ConductivityMolarconductivity()
But if solution contains 1 gm mole of the
electrolyte therefore, the measured conductance will be
the molar conductivity. Thus,
Molar conductivity() 100 Conductivity
In other words,( ) V
where V is the volume of the solution in 3 cm containing one gram mole of the electrolyte.
If M is the concentration of the solution in mole per
litre, then
M mole of electrolyte is present in
3 1000 cm
1 mole of electrolyte is present in
(^1000 ) cm M
of
solution
Thus,
3 Volume^ in cm containing 1 mole of
electrolyte.
or M
1000
Units of Molar Conductance : The units of molar
conductance can be derived from the formula ,
M
1000
The units of are S
1 cm and units of are,
2 1 2 1
3 1 Scm mol S cm mol mol
cm Λ S cm
According to SI system, molar conductance is
expressed as ,
2 1 Sm mol if concentration is expressed as
.
3 mol m
(7) Equivalent conductivity : It is defined as the
conducting power of all the ions produced by dissolving
one gram equivalent of an electrolyte in solution.
It is expressed as e and is related to specific
conductance as
C M
e
1000 1000
( M is Molarity of the
solution)
where C is the concentration in gram equivalent
per litre (or Normality). This term has earlier been
quite frequently used. Now it is replaced by molar
conductance. The units of equivalent conductance are
( ).
1 2 1 Ohm cm gmequiv
(8) Experimental measurement of conductance
(i) The conductance of a solution is reciprocal of
the resistance, therefore, the experimental
determination of the conductance of a solution involves
the measurement of its resistance.
(ii) Calculation of conductivity : We have seen
that conductivity ( ) is reciprocal of resistivity ( ), i.e.,
and l
a R
a
l G a
l
R
or
1
where G is the conductance of the cell, l is the
distance of separation of two electrodes having cross
section area.
2 a cm
The quantity
a
l is called cell constant and is
expressed in.
1 ^ cm Knowing the value of cell constant
and conductance of the solution, the specific
conductance can be calculated as,
G Cellconstant
i.e., Conductivi tyConductanceCellconstant
Factors affecting the electrolytic conductance
In general, conductance of an electrolyte depends
upon the following factors,
(1) Nature of electrolyte : The conductance of an
electrolyte depends upon the number of ions present in
the solution. Therefore, the greater the number of ions
in the solution the greater is the conductance. The
number of ions produced by an electrolyte depends
upon its nature. The strong electrolytes dissociate
almost completely into ions in solutions and, therefore,
their solutions have high conductance. On the other
hand, weak electrolytes, dissociate to only small
extents and give lesser number of ions. Therefore, the
solutions of weak electrolytes have low conductance.
(2) Concentration of the solution : The molar
conductance of electrolytic solution varies with the
concentration of the electrolyte. In general, the molar
conductance of an electrolyte increases with decrease
in concentration or increase in dilution.
The molar conductance of strong electrolyte
( HCl , KCl , KNO 3 ) as well as weak electrolytes
3 4 CHCOOHNHOH increase with decrease in
concentration or increase in dilution. The variation is
however different for strong and weak electrolytes.
The variation of molar conductance with
concentration can be explained on the basis of
conducting ability of ions for weak and strong
electrolytes.
For weak electrolytes the variation of with
dilution can be explained on the bases of number of
ions in solution. The number of ions furnished by an
expressing the molar conductivity of an electrolyte is
illustrated as,
The molar conductivity of HCl at infinite dilution
can be expressed as,
(^) HCl H H Cl Cl ; For HCl, H 1 and
^1. Cl
So, ( 1 ) ( 1 ) (^) HCl H Cl ; Hence,
(^) HCl H Cl
(2) Applications of Kohlrausch's law : Some
typical applications of the Kohlrausch's law are described
below,
(i) Determination of
(^) m for weak electrolytes :
The molar conductivity of a weak electrolyte at infinite
dilution ( ) m cannot be determined by extrapolation
method. However, (^) m values for weak electrolytes can
be determined by using the Kohlrausch's equation.
(^) CH COOH CHCOONa HCl NaCl 3 3
(ii) Determination of the degree of ionisation of
a weak electrolyte : The Kohlrausch's law can be used
for determining the degree of ionisation of a weak
electrolyte at any concentration. If
c m is the molar
conductivity of a weak electrolyte at any concentration
C and,
m is the molar conductivity of a electrolyte at
infinite dilution. Then, the degree of ionisation is given
by, ( )
c m
m
c m c
Thus, knowing the value of
c m , and
m (From
the Kohlrausch's equation), the degree of ionisation at
any concentration( ) c can be determined.
(iii) Determination of the ionisation constant of
a weak electrolyte : Weak electrolytes in aqueous
solutions ionise to a very small extent. The extent of
ionisation is described in terms of the degree of
ionisation ( ).In solution, the ions are in dynamic
equilibrium with the unionised molecules. Such an
equilibrium can be described by a constant called
ionisation constant. For example, for a weak
electrolyte AB , the ionisation equilibrium is, AB ⇌
A B ; If C is the initial concentration of the
electrolyte AB in solution, then the equilibrium
concentrations of various species in the solution are,
[ AB ] C ( 1 ), A C
[ ] and B C [ ]
Then, the ionisation constant of AB is given by,
( 1 ) ( 1 )
.
[ ]
[ ][ ]
2
C
C
C C
AB
A B K
We know, that at any concentration C, the degree
of ionisation ( )is given by,
m
c m /
Then, ( )
( )
[ 1 ( / )]
( / )
2 2
c m m m
c m
m
c m
m
c C m C K
; Thus,
knowing
m and
c m at any concentration, the
ionisation constant ( K ) of the electrolyte can be
determined.
(iv) Determination of the solubility of a
sparingly soluble salt : The solubility of a sparingly
soluble salt in a solvent is quite low. Even a saturated
solution of such a salt is so dilute that it can be
assumed to be at infinite dilution. Then, the molar
conductivity of a sparingly soluble salt at infinite
dilution( )
m can be obtained from the relationship,
m ........(i)
The conductivity of the saturated solution of the
sparingly soluble salt is measured. From this, the
conductivity of the salt ( (^) salt )can be obtained by using
the relationship, salt sol wate r , where, water is the
conductivity of the water used in the preparation of the
saturated solution of the salt.
Cm
salt salt
........(ii)
From equation (i) and (ii) ;
(^1000) salt
C m , Cm is the molar concentration
of the sparingly soluble salt in its saturated solution.
Thus, Cm is equal to the solubility of the sparingly
soluble salt in the mole per litre units. The solubility of
the salt in gram per litre units can be obtained by
multiplying Cm with the molar mass of the salt.
Electrochemical or Galvanic cell
“ Electrochemical cell or Galvanic cell is a device in
which a spontaneous redox reaction is used to convert
chemical energy into electrical energy i.e. electricity can
be obtained with the help of oxidation and reduction
reaction ”.
(1) Characteristics of electrochemical cell :
Following are the important characteristics of
electrochemical cell,
Voltmeter (^) Salt bridge
Porous plug
Zn anode
Cu cathode
e
- e -
ZnSO 4^ CuSO 4
Fig. 12.
(i) Electrochemical cell consists of two vessels,
two electrodes, two electrolytic solutions and a salt
bridge.
(ii) The two electrodes taken are made of
different materials and usually set up in two separate
vessels.
(iii) The electrolytes are taken in the two
different vessels called as half - cells.
(iv) The two vessels are connected by a salt
bridge/porous pot.
(v) The electrode on which oxidation takes place
is called the anode (or – ve pole) and the electrode on
which reduction takes place is called the cathode (or +
ve pole).
(vi) In electrochemical cell, ions are discharged
only on the cathode.
(vii) Like electrolytic cell, in electrochemical cell,
from outside the electrolytes electrons flow from anode
to cathode and current flow from cathode to anode.
(viii) For electrochemical cell,
Ecell ve , G ve.
(ix) In a electrochemical cell, cell reaction is
exothermic.
(2) Salt bridge and its significance
(i) Salt bridge is U – shaped glass tube filled with
a gelly like substance, agar – agar (plant gel) mixed
with an electrolyte like KCl, KNO 3 , NH 4 NO 3 etc.
(ii) The electrolytes of the two half-cells should
be inert and should not react chemically with each
other.
(iii) The cation as well as anion of the electrolyte
should have same ionic mobility and almost same
transport number, viz. KCl , KNO 3 , NH 4 NO 3 etc.
(iv) The following are the functions of the salt
bridge,
(a) It connects the solutions of two half - cells
and completes the cell circuit.
(b) It prevent transference or diffusion of the
solutions from one half cell to the other.
(c) It keeps the solution of two half - cells
electrically neutral.
(d) It prevents liquid – liquid junction potential
i.e. the potential difference which arises between two
solutions when they contact with each other.
(3) Representation of an electrochemical cell
The cell may be written by arranging each of the
pair left – right, anode – cathode, oxidation – reduction,
negative and positive in the alphabetical order as,
(4) Reversible and irreversible cells : A cell is said
to be reversible if the following two conditions are
fulfilled
(i) The chemical reaction of the cell stops when
an exactly equal external emf is applied.
(ii) The chemical reaction of the cell is reversed
and the current flows in opposite direction when the
external emf is slightly higher than that of the cell. Any
other cell, which does not obey the above two
conditions, is termed as irreversible. Daniell cell is
reversible but Zn | HSO | Ag 2 4 cell is irreversible in
nature
(5) Types of electrochemical cells : Two main
types of electrochemical cells have been reported, these
are,
(i) Chemical cells : The cells in which electrical
energy is produced from the energy change
accompanying a chemical reaction or a physical process
are known as chemical cells. Chemical cells are of two
types,
(a) Chemical cells without transference : In this
type of chemical cells, the liquid junction potential is
neglected or the transference number is not taken into
consideration. In these cells, one electrode is reversible
to cations while the other is reversible to the anions of
the electrolyte.
(b) Chemical cells with transference : In this type
of chemical cells, the liquid-liquid junction potential or
diffusion potential is developed across the boundary
between the two solutions. This potential develops due
to the difference in mobilities of ve and ve ions of
the electrolytes.
(6) Concentration cells : “ A cell in which
electrical energy is produced by the transference of a
substance from a system of high concentration to one at
low concentration is known as concentration cells”.
Concentration cells are of two types.
(i) Electrode concentration cells : In these cells,
the potential difference is developed between two
electrodes at different concentrations dipped in the
same solution of the electrolyte. For example, two
hydrogen electrodes at different gaseous pressures in
the same solution of hydrogen ions constitute a cell of
this type.
Cathode
(pressure ) | | Anode
, 2 ( pressure 1 ) H 2 p 2 Pt H
Pt H p ;
log 2
2
1 cell p
p E at C
o 25 If p 1 (^) p 2 , oxidation occurs
at L. H. S. electrode and reduction occurs at R. H. S.
electrode.
In the amalgam cells, two amalgams of the same
metal at two different concentrations are immersed in
the same electrolytic solution. M ( HgC 1 )| M | Zn ( HgC 2 )
n
Left
Anode
Oxidation
Negative
Bridge Right
Cathode
Reductio
n Positive
Cathode : Graphite rod Anode : Zn pot
Electrolyte : Paste of NH 4 (^) Cl ZnCl 2 in starch
Emf : 1.2 V to 1.5 V
At cathode : NH (^) 4 MnO 2 2 e MnO ( OH ) NH 3
At Anode : Zn Zn 2 e
2
Over all reaction :
3
2 Zn NH 4 MnO 2 Zn MnO ( OH ) NH
(iv) Mercury cell
Cathode : Mercury (II) oxide Anode :
Zn rod
Electrolyte : Paste of KOH ZnO Emf :
1.35 V
At cathode :
HgO ( s ) H 2 O ( l ) 2 e Hg ( l ) 2 OH ( aq )
At Anode :
Zn (^) s 20 H ( aq ) ZnO ( s ) H 2 O ( l ) 2 e (amalgam)
()
Over all reaction : Zn ( (^) s ) HgO ( s ) ZnO ( s ) Hg ( l )
(2) Secondary cells : In the secondary cells, the
reactions can be reversed by an external electrical
energy source. Therefore, these cells can be recharged
by passing electric current and used again and again.
These are also celled storage cells. Examples of
secondary cells are, lead storage battery and nickel –
cadmium storage cell.
In charged Lead storage cell Alkali cell
Positive
electrode
Perforated lead plates coated with PbO 2 Perforated steel plate coated with Ni ( OH ) 4
Negative
electrode
Perforated lead plates coated with pure lead Perforated steel plate coated with Fe
Electrolyte dil. H 2 SO 4 20% solution of KOH + 1% LiOH
During charging Chemical reaction
At anode : PbSO 4 + 2 H
At cathode : PbSO 4 + SO 4
PbO 2
2 H 2 SO 4
Specific gravity of H 2 SO 4 increases and when
specific gravity becomes 1.25 the cell is fully
charged.
Emf of cell : When cell is fully charged then E =
2.2 volt
Chemical reaction
At anode : Ni ( OH ) 2 + 2 OH
Ni ( OH ) 4
At cathode : Fe ( OH ) 2 + 2 K
2 KOH
Emf of cell : When cell is fully charged
then E = 1.36 volt
During
discharging
Chemical reaction
At anode : Pb + SO 4
At cathode : PbO 2 + 2 H
2 H 2 O
Specific gravity of H 2 SO 4 decreases and when
specific gravity falls below 1.18 the cell
requires recharging.
Chemical reaction
At anode : Fe + 2 OH
At cathode : Ni ( OH ) 4 + 2 K
Ni ( OH ) 2 +
2 KOH
Emf of cell : When emf of cell falls below
- 1 V it requires charging.
Glass
vessel
PbO 2
Pb
dil. H 2 SO 4
Ni ( OH ) 2
Perforated
steel grid
KOH 20%
+^ –
Fe ( OH ) 2
Emf of cell : When emf of cell falls below 1.
volt the cell requires recharging.
Efficiency 80% 60%
Fuel cells
These are Voltaic cells in which the reactants are
continuously supplied to the electrodes. These are
designed to convert the energy from the combustion of
fuels such as H 2 , CO , CH 4 , etc. directly into electrical
energy. The common example is hydrogen-oxygen fuel
cell as described below,
In this cell, hydrogen and oxygen are bubbled
through a porous carbon electrode into concentrated
aqueous sodium hydroxide or potassium hydroxide.
Hydrogen (the fuel) is fed into the anode compartment
where it is oxidised. The oxygen is fed into cathode
compartment where it is reduced. The diffusion rates of
the gases into the cell are carefully regulated to get
maximum efficiency. The net reaction is the same as
burning of hydrogen and oxygen to form water. The
reactions are
At anode :
2 H (^) 2 ( g ) 2 OH 2 H 2 O ( l ) 2 e
At cathode : O 2 ( g ) 2 H 2 O ( l ) 4 e 4 OH ( aq )
Overall reaction :
2 H 2 (^) ( g ) O 2 ( g ) 2 H 2 O ( l )
Each electrode is made of porous compressed
carbon containing a small amount of catalyst
( Pt , Ag or CoO ). This cell runs continuously as long as
the reactants are fed. Fuel cells convert the energy of
the fuel directly into electricity EMF of fuel cell is 1.
V. This cell has been used for electric power in the
Apollo space programme. The important advantages of
fuel cells are
(1) High efficiency : The fuel cells convert the
energy of a fuel directly into electricity and therefore,
they are more efficient than the conventional methods
of generating electricity on a large scale by burning
hydrogen, carbon fuels. Though we expect 100 %
efficiency in fuel cells, so far 60 – 70% efficiency has
been attained. The conventional methods of production
of electrical energy involve combustion of a fuel to
liberate heat which is then used to produce electricity.
The efficiency of these methods is only about 40%.
(2) Continuous source of energy : There is no
electrode material to be replaced as in ordinary
battery. The fuel can be fed continuously to produce
power. For this reason, H (^) 2 O 2 fuel cells have been used
in space crafts.
(3) Pollution free working : There are no
objectionable byproducts and, therefore, they do not
cause pollution problems. Since fuel cells are efficient
and free from pollution, attempts are being made to get
better commercially practical fuel cells.
Electrode Potential
(1) When a metal ( M ) is placed in a solution of its
ions ( M ++ ), either of the following three possibilities
can occurs, according to the electrode potential
solution pressure theory of Nernst.
(i) A metal ion M n + collides with the electrode,
and undergoes no change.
(ii) A metal ion M n + collides with the electrode,
gains n electrons and gets converted into a metal atom
M, ( i.e. the metal ion is reduced).
M ( aq ) ne M ( s )
n
(iii) A metal atom on the electrode M may lose an
electrons to the electrode, and enter to the solution as
n M , ( i.e. the metal atom is oxidised).
M s M aq ne
n () ( ).
Thus, “ the electrode potential is the tendency of an
electrode to lose or gain electrons when it is in contact
with solution of its own ions .”
(2) The magnitude of electrode potential depends
on the following factors,
(i) Nature of the electrode, (ii) Concentration of
the ions in solution, (iii) Temperature.
(3) Types of electrode potential : Depending on
the nature of the metal electrode to lose or gain
electrons, the electrode potential may be of two types,
(i) Oxidation potential : When electrode is
negatively charged with respect to solution, i.e., it acts
as anode. Oxidation occurs.
M M ne
n
(ii) Reduction potential : When electrode is
positively charged with respect to solution, i.e. it acts
as cathode. Reduction occurs. M ne M n
(4) Standard electrode potential : “ If in the half
cell, the metal rod (M) is suspended in a solution of one
molar concentration, and the temperature is kept at 298
K, the electrode potential is called standard electrode
potential, represented usually by
o E ”. ‘or’
The standard electrode potential of a metal may
be defined as “ the potential difference in volts developed
in a cell consisting of two electrodes, the pure metal in
H 2
Anode– + Cathode
O 2
Electrolyte
OH
H 2 O
Fig. 12.
Porous carbon electrode
It is responsible for the
steady flow of current in
the cell.
It is not responsible for
the steady flow of
current in the cell.
(4) Cell EMF and the spontaneity of the reaction :
We know, G nFEcell
Nature of
reaction
G(or G)
o Δ Δ E (orE )
o cell cell
Spontaneous – +
Equilibrium 0 0
Non – spontaneous + –
Nernst's equation
(1) Nernst’s equation for electrode potential
The potential of the electrode at which the
reaction,
M ( aq ) ne M ( s )
n
takes place is described by the equation,
[ ( .)]
[ ()]
ln
0 / / M aq
Ms
nF
RT
E E
M n^ M Mn M n
^ ^
or [ ( )]
[ ()]
log
0 2.^303
/ / M aq
Ms
nF
RT
E E
M n^ M Mn M n
^ ^
above eq. is called the Nernst equation.
Where,
M M
E (^) n / ^ =^ the potential of the electrode at a given
concentration,
0 M / M
E (^) n = the standard electrode potential
R = the universal gas constant,
1 1
- 31
JK mol
T = the temperature on the absolute scale,
n = the number of electrons involved in the
electrode reaction,
F = the Faraday constant : (96500 C ),
[ M ( s )] = the concentration of the deposited metal,
[ M ( aq )]
n = the molar concentration of the metal
ion in the solution,
The concentration of pure metal M ( s ) is taken as
unity. So, the Nernst equation for the M M
n /
electrode is written as,
[ ( )]
log
0 2.^303
/ / nF M aq
RT
E E
M n^ M Mn M n
^ ^
At 298 K, the Nernst equation for the M M n /
electrode can be written as,
[ ( )]
log
0 0.^0591
/ / n M aq
E E
M n^ M Mn M n
^ ^
For an electrode (half - cell) corresponding to the
electrode reaction,
Oxidised form (^) ne Reduced form
The Nernst equation for the electrode is written
as,
[Oxidisedform]
[Reducedform] log
0 2.^303
nF
RT
E E
half cell halfcell
At 298 K, the Nernst equation can be written as,
[Oxidisedform]
[Reducedform] log
0 0.^0591
n
E (^) half cell Ehalf cell
(2) Nernst’s equation for cell EMF
For a cell in which the net cell reaction involving
n electrons is, aA bB cC dD
The Nernst equation is written as,
a b
d
cell cell A B
D
nF
RT E E [][]
[C][ ] ln
c 0
Where,
0 0 0 Ecell Ecathode E anode.
The o Ecell is called the standard cell potential.
or a b
c d o cell A B
C D
nF
RT E E [][]
[ ][ ] log
- 303 cell
At 298 K, above eq. can be written as,
or a b
c d o cell A B
C D
n
E E [][ ]
[ ][ ] log
- 0592 cell
It may be noted here, that the concentrations of A,
B, C and D referred in the eqs. are the concentrations at
the time the cell emf is measured.
(3) Nernst’s equation for Daniells cell :
Daniell’s cell consists of zinc and copper electrodes.
The electrode reactions in Daniell’s cell are,
At anode :
Zn ( s ) Zn ( aq ) 2 e
2
At cathode : ( ) 2 ()
2 Cu aq e Cus
Net cell reaction :
2 2 Zn s Cu aq Cus Zn aq
Therefore, the Nernst equation for the Daniell’s
cell is,
[ ()][ ( )]
[ ()][ ( )]
log 2
2
2 0
Zns Cu aq
Cus Zn aq
F
RT
E (^) cdll Ecell
Since, the activities of pure copper and zinc
metals are taken as unity, hence the Nernst equation
for the Daniell’s cell is,
[ ( )]
[ ( ]
log 2
2
2 0
Cu aq
Zn aq
F
RT
E (^) cdll Ecell
The above eq. at 298 K is,
V
Cu aq
Zn aq E E
o cdll cell [ ( )]
[ ( ]
log 2
2
2
For Daniells cell, Ecell 1. 1 V
0
(4) Nernst's equation and equilibrium constant
For a cell, in which the net cell reaction involving
n electrons is, aA bB cC dD
The Nernst equation is
a b
c d
Cell cell A B
C D
nF
RT E E [][ ]
[ ][ ] ln 0 .....(i)
At equilibrium, the cell cannot perform any useful
work. So at equilibrium, Cell E is zero. Also at
equilibrium, the ratio
c equilibrium
a b
c d
a b
c d K A B
C D
A B
C D
[ ][]
[ ][ ]
[][ ]
[ ][ ]
Relationship between potential, Gibbs energy
and equilibrium constant
The electrical work (electrical energy) is equal to
the product of the EMF of the cell and electrical charge
that flows through the external circuit i.e. ,
W max nFEcell ......(i)
According to thermodynamics the free energy
change ( G )is equal to the maximum work. In the cell
work is done on the surroundings by which electrical
energy flows through the external circuit, So
W (^) max, G
......(ii)
from eq. (i) and (ii) G nFEcell
In standard conditions
0 0 cell G nFE
Where
0 G standard free energy change
But cell RT Kc nF
E log
0 2.^303
RT Kc
nF
G nF log 0 2.^303
2.303 RTlogKc
0 G or G G 2. 303 RT log Q
ln ( 2. 303 log ln )
0 G RT Kc X X
Electrochemical series
(1) The standard reduction potentials of a large
number of electrodes have been measured using
standard hydrogen electrode as the reference electrode.
These various electrodes can be arranged in increasing
or decreasing order of their reduction potentials. The
arrangement of elements in order of increasing
reduction potential values is called electrochemical
series .It is also called activity series , of some typical
electrodes.
(2) Characteristics of Electrochemical series
(i) The negative sign of standard reduction
potential indicates that an electrode when joined with
SHE acts as anode and oxidation occurs on this
electrode. For example, standard reduction potential of
zinc is – 0.76 volt , When zinc electrode is joined with
SHE, it acts as anode (– ve electrode) i.e. , oxidation
occurs on this electrode. Similarly, the + ve sign of
standard reduction potential indicates that the
electrode when joined with SHE acts as cathode and
reduction occurs on this electrode.
(ii) The substances, which are stronger reducing
agents than hydrogen are placed above hydrogen in the
series and have negative values of standard reduction
potentials. All those substances which have positive
values of reduction potentials and placed below
hydrogen in the series are weaker reducing agents than
hydrogen.
(iii) The substances, which are stronger oxidising
agents than
^ H ion are placed below hydrogen in the
series.
(iv) The metals on the top (having high negative
value of standard reduction potentials) have the
tendency to lose electrons readily. These are active
metals. The activity of metals decreases from top to
bottom. The non-metals on the bottom (having high
positive values of standard reduction potentials) have
the tendency to accept electrons readily. These are
active non-metals. The activity of non-metals increases
from top to bottom.
Table : 12.3 Standard reduction electrode potentials at
298K
Element Electrode Reaction
(Reduction)
Standard
Electrode
Reduction
potential E 0 ,
volt
Li Li
_K K
- = K – 2.
Ba Ba ++ + 2 e = Ba – 2.
Sr Sr ++ + 2 e = Sr – 2.
Ca Ca e Ca 2 2 – 2.
Na Na e Na –2.
Mg (^) Mg^2 ^ 2 e Mg – 2.
Al (^) Al^3 ^ 3 e Al – 1.
tendency to accept electrons Increasing^ strength as oxidising agent rength as reducing agent Increasing^ ndency to lose electro
ns
The metals which are below hydrogen in
electrochemical series like Cu, Hg, Au, Pt, etc., do not
evolve hydrogen from dilute acids.
(d) Displacement of hydrogen from water : Iron
and the metals above iron are capable of liberating
hydrogen from water. The tendency decreases from top
to bottom in electrochemical series. Alkali and alkaline
earth metals liberate hydrogen from cold water but Mg ,
Zn and Fe liberate hydrogen from hot water or steam.
(iv) Reducing power of metals : Reducing nature
depends on the tendency of losing electron or electrons.
More the negative reduction potential, more is the
tendency to lose electron or electrons. Thus reducing
nature decreases from top to bottom in the
electrochemical series. The power of the reducing agent
increases, as the standard reduction potential becomes
more and more negative. Sodium is a stronger reducing
agent than zinc and zinc is a stronger reducing agent
than iron. (decreasing order of reducing nature)
Na Zn Fe
Reduction potential
Element
Alkali and alkaline earth metals are strong
reducing agents.
(v) Oxidising nature of non-metals : Oxidising
nature depends on the tendency to accept electron or
electrons. More the value of reduction potential, higher
is the tendency to accept electron or electrons. Thus,
oxidising nature increases from top to bottom in the
electrochemical series. The strength of an oxidising
agent increases as the value of reduction potential
becomes more and more positive.
F 2 (Fluorine) is a stronger oxidant than Cl 2 , Br 2
and I 2 , Cl 2 (Chlorine) is a stronger oxidant than
Br 2 and I 2
Oxidisingnatureincreases
2 2 2 2
I Br Cl F
Thus, in electrochemical series
(vi) Thermal stability of metallic oxides : The
thermal stability of the metal oxide depends on its
electropositive nature. As the electropositivity
decreases from top to bottom, the thermal stability of
the oxide also decreases from top to bottom. The oxides
of metals having high positive reduction potentials are
not stable towards heat. The metals which come below
copper form unstable oxides, i.e., these are decomposed
on heating.
2 2 2
1 Ag O 2 Ag O
2 HgO 2 Hg O 2
; Nodecomposition
2 3
2
Al O
NaO
BaO
(vii) Extraction of metals : A more
electropositive metal can displace a less electropositive
metal from its salt's solution. This principle is applied
for the extraction of Ag and Au by cyanide process.
silver from the solution containing sodium argento
cyanide, NaAg ( CN ) 2 , can be obtained by the addition of
zinc as it is more electro-positive than Ag.
2 NaAg ( CN ) Zn NaZn ( CN ) 2 Ag 2 2 4
Corrosion
(1) When metals are exposed to atmospheric
conditions, they react with air or water in the
environment to form undesirable compounds (usually
oxides). This process is called corrosion. Almost all
metals except the least active metals such as gold,
platinum and palladium are attacked by environment
i.e ., undergo corrosion. For example, silver tarnishes,
copper develops a green coating, lead or stainless steel
lose their lusture due to corrosion. Corrosion causes
enormous damage to building, bridges, ships and many
other articles made of iron.
Thus corrosion is a process of deterioration of a
metal as a result of its reaction with air or water
(environment) surrounding it.
In case of iron, corrosion is called rusting.
Chemically, rust is hydrated form of ferric oxide,
2 3 FeO. xH O 2
. Rusting of iron is generally caused by
moisture, carbon dioxide and oxygen present in air. It
has been observed that rusting takes place only when
iron is in contact with moist air. Iron does not rust in
dry air and in vacuum.
(2) Factors which affect corrosion : The main
factors which affect corrosion are
More the reactivity of metal, the more will be the
possibility of the metal getting corroded.
The impurities help in setting up voltaic cells,
which increase the speed of corrosion
Presence of electrolytes in water also increases
the rate of corrosion
Oxidising
nature
nature Reducing
Top
Botto
m
(Strongest reducing agent)
Highest negative reduction potential
or
(Minimum reduction potential)
(Strongest oxidising agent)
Highest positive value of reduction potential
Element :
Reduction potential :
Presence of CO 2 in natural water increase rusting
of iron.
(v) When the iron surface is coated with layers of
metals more active than iron, then the rate of corrosion
is retarded.
A rise in temperature (with in a reasonable limit)
increases the rate of corrosion.
(3) Classification of corrosion process :
Depending upon the nature of corrosion, and the
factors affecting it, the corrosion may be classified as
follows.
(i) Chemical corrosion : Such corrosion, generally
takes place when
(a) Reactive gases come in contact with metals at
high temperatures e.g ., corrosion in chemical industry.
(b) Slow dissolution of metal takes place when
kept in contact with non conducting media containing
organic acids.
(ii) Bio-chemical corrosion or Bio-corrosion :
This is caused by the action of microorganisms. Soils of
definite composition, stagnant water and certain
organic products greatly favour the bio-corrosion.
(iii) Electrochemical corrosion : It occurs in a
gaseous atmosphere in the presence of moisture, in
soils and in solutions.
(4) Mechanism of rusting of iron :
Electrochemical theory of rusting.
The overall rusting involves the following steps,
(i) Oxidation occurs at the anodes of each
electrochemical cell. Therefore, at each anode neutral
iron atoms are oxidised to ferrous ions.
At anode : ( ) ( ) 2.
2 Fes Fe aq e
Thus, the metal atoms in the lattice pass into the
solution as ions, leaving electrons on the metal itself.
These electrons move towards the cathode region
through the metal.
(ii) At the cathodes of each cell, the electrons are
taken up by hydrogen ions (reduction takes place). The
H ions are obtained either from water or from acidic
substances (e.g. CO 2 )in water
H 2 O H OH or
CO 2 H 2 O H HCO 3
At cathode : H e H
The hydrogen atoms on the iron surface reduce
dissolved oxygen. 4 H O 2 2 H 2 O
Therefore, the overall reaction at cathode of
different electrochemical cells may be written as,
4 H O 2 4 e 2 H 2 O
(iii) The overall redox reaction may be written by
multiplying reaction at anode by 2 and adding reaction
at cathode to equalise number of electrons lost and
gained i.e.
Oxi. half reaction : ( ) ( ) 2 ] 2 2 Fes Fe aq e
( E 0. 44 V )
Red. half reaction : 4 H O 2 4 e 2 H 2 O
( E 1. 23 V )
Overall cell reaction : Fe s H O Fe aq H 2 O
2 2 () 4 2 2 ( ) 2
( 1. 67 ) Cell E V
The ferrous ions are oxidised further by
atmospheric oxygen to form rust.
4 Fe ( aq ) O 2 ( g ) 4 H 2 O 2 Fe 2 O 3 8 H 2 and
Rust
2 3 2 2 3 2 Fe O xHO FeO. xHO
It may be noted that salt water accelerates
corrosion. This is mainly due to the fact that salt water
increases the electrical conduction of electrolyte
solution formed on the metal surface. Therefore,
rusting becomes more serious problem where salt
water is present.
(5) Corrosion protection : Corrosion of metals
can be prevented in many ways. Some commonly used
methods are
(i) By surface coating
(a) By applying, oil, grease, paint or varnish on
the surface.
(b) By coating/depositing a thin layer of any other
metal which does not corrode. For example, iron
surface can be protected from corrosion by depositing a
thin layer of zinc, nickel or chromium on it.
Copper/brass can be protected by coating it with a thin
layer of tin. Tinning of brass utensils is a very common
practice in our country.
(c) By Galvanization : Prevention of corrosion of
iron by Zn coating.
(ii) By connecting metal to a more
electropositive metal : As long as the more
electropositive metal is there, the given metal does not
get corroded. For example, iron can be protected from
corrosion by connecting it to a block/plate of zinc or
Fe 2+
2e–
+
Flow of electron s
Fe anode
Iron
Rust
Drop of moisture 4 H ++O 2 +4 e – 2 H 2 O (Cathode)
Schematic representation of mechanism of rusting of iron Fig. 12.
