Module 1 electrode potential & cells - battery technology

ELECTRODE POTENTIAL & CELLS
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ENGG.CHEMISTRY [RM] Page 1
MODULE-1
Electrochemistry: It is a branch of chemistry which deals with the study of transformation of chemical
energy into electrical energy and vice versa”
Electrochemical cell and Classification with examples:
An electrochemical cell is a device, which is used to convert chemical energy into electrical energy and vice
versa.
These electrochemical cells are classified into two types as follows.
1) Galvanic or Voltaic cells: These are the electrochemical cells, which convert chemical energy into electrical
energy.
Ex. Daniel cell, Dry cell, etc
2) Electrolytic cells-are devices which convert electrical energy into chemical energy.
Example: Electrolysis of molten NaCl, Recharge process of lead acid battery
Galvanic or Voltaic cells:
Galvanic or Voltaic cells are again classified into three types as follows
a) Primary cells: These are the cells which serve as a source of energy only as long as the active chemical
species are present in the cell. The cell reactions are irreversible. These are designed for only single discharge
and cannot be charged again.
Ex: Dry Cell, Zn – Hgo cell, Zn-Ag2o cell etc.
b) Secondary cells: These cells are chargeable and can be used again and again. The cell reactions are
reversible and are often called as reversible cells. During discharging the cells acts like voltaic cell converting
chemical energy into electrical energy. During charging the cell acts like electrolytic cell by converting electric
energy into chemical energy, hence these batteries are called as storage battery.
Ex: Lead acid Battery, Ni-cd cells. Lithium ion cells etc.
c) Concentration of cells:
These are the electrochemical cells consisting of same metalelectrodes dipped in same metal ionic solution in
both the half cells butare different in the concentration of the metal ions.
Ex: Cu/Cu2+
(M1) || Cu2+
(M2)/Cu
Ex: Copper concentration cell, Zinc concentration cell
Oxidation: A species loses one or more electrons resulting in the increase in its oxidation number.
Reduction: A species gain one or more electrons resulting in a decreasing in its oxidation number.
Oxidation should accompanied by reduction, because if one losses electrons another must ready to accept
electrons. This reaction is called redox reaction.
Single electrode Potential:
It is defined as the potential developed at the interphase between the metal and the solution, when a
metal is dipped in a solution containing its own ions. It is represented as E

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MODULE-1
Standard reduction potential (Eo
) :
It is defined as potential developed at the interface between the metal and the solution, when a metal is dipped
in a solution containing its own ions of unit concentration at 298K. [If the electrodes involve gases then it is one
atmospheric pressure] It is denoted as E0
.
Electromotive force (EMF):
It is defined as the potential difference between the two electrodes of a galvanic cell which causes the flow of
current from an electrode with higher reduction potential to the electrode with lower reduction potential.
It is denoted as E cell.
E cell = E right –E left.
E cell = E cathode – E anode.
Electrochemical series:
The arrangement of elements in the order of their standard reduction potential is reffered to as emf or
electrochemical series. Such a arrangement of few elements given in the table.
Mn+
/M Eo
(volts) Mn+
/M Eo
(volts)
Li+
/Li -3.05 H+
/H2 0.00
K+
/K -2.93 Sn4+
/ Sn2+
+0.15
Mg+
/Mg -2.37 Cu2+
/Cu +0.16
Al3+
/Al -1.66 Cu+
/Cu +0.52
Zn2+
/Zn -0.76 I2/I-
+0.54
Cr3+
/Cr -0.74 Fe3+
/Fe2+
+0.77
Fe2+
/Fe -0.44 Hg2+
/Hg+
+0.79
Cr3+
/Cr2+
-0.41 Ag+
/Ag +0.80
Cd2+
/Cd -0.40 Hg2+
/Hg +0.85
Ni2+
/Ni -0.25 Pt2+
/Pt +1.20
Sn2+
/Sn -0.14 Cr7+
/Cr3+
+1.31
Pb2+
/Pb -0.13 Cl2/2Cl-
+1.36
Fe3+
/Fe2+
-0.041 Au3+
/Au +1.50
1) A negative value indicates oxidation tendency and a positive value indicates reduction tendency with respect
to hydrogen.
2) The metal with lower electrode potential is more reactive and as the electrode potential increases, the
reactivity decreases, and metals with higher electrode potentials are nobler.
3) Metals with lower electrode potentials have the tendency to replace metals with higher electrode potential
from their solutions for example, Zn displaces Cu, and Cu displaces Ag
4) Metals with negative electrode potentials can liberate hydrogen from acidic solutions
Derivation of Nernst Equation for Electrode potential:
In 1889 Nernst derived a quantitative relationship between the electrode potential and the concentrations of
metal ions are involved. The maximum work available from a reversible chemical process is equal to the
maximum amount of electrical energy that can be obtained; it shows decrease in free energy.
Wmax = – ∆G------------------------------------[1]

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MODULE-1
And
Wmax = difference in potential between two electrode x total quantity of charge flowing through the cell
Total quantity of charge flowing through the cell = (No. of moles of electrons) x (Faradays constant)
So Wmax = nFEcell -----------------------------[2]
Where,
E = Electrode potential
E0
= standard electrode potential
n = no. of electrons
[Mn+
] = Concentration of metal ions
R = Universal gas constant = 8.314J K-1
mol-1
T = Temperature (In Kelvin) = 298K
]7[]n[Mlog
0591.00 
n
E
cell
E
]5[
c
Kln0GG
,isotherm'reactionhoffvant'abyrelatedare0GandG,
c
K
]4[
]n[M
[M]
c
K
aswrittenbecan
c
Kconstantequlibrium,reactionabovefor the
M-nenM
reaction,electrodereversibleaconsider
0nFE-0Gion,std.conditunder
[3]-----nFE-G
[2]&[1]eqnequate








RT
]6[]nlog[M
303.20EE
1[M]condition,standardunder
]n[M
[M]
lnRT0EE
nF-bysidesboth thedivide
]n[M
[M]
ln0-nFEnFE-
equation,abovetheto0GandG,
c
Kofvaluesthesubstitute







nF
RT
nF
RT
RT

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MODULE-1
 
 reactant
product
cKwherelog
2
0591.00 
c
KE
cell
E
 
 anodeatSpecies
cathodeatSpecies
log
2
0591.00  E
cell
E
Concentrations cells:
“A concentration cell is an electrochemical cell in which electrode materials and electrolytes of two half
cells are composed of same material but the concentration of two solutions are different”
Ex- Cu/Cu2+
(M1)|| Cu2+
(M2)/Cu
A concentration cell consists of two same metal electrode dipped into their own ionic solutions of two different
concentration.
Thus in a concentration cell, the electrode with lower electrolyte concentration acts as anode and the one
with higher electrolyte concentration acts as cathode. The concentration of ions at anode increases and at
cathode decreases, when the cell is in operation.
Consider two copper rods are dipped into their own ionic solutions of M1 and M2 and it is represented as
Cu/Cu2+
(M1) Cu2+
(M2)/Cu
By electrochemical conventions, if M2 > M1 then, we have the following reactions.
At anode



 eMCu
s
Cu 2)1(
2
)(
At cathode
)(
2)2(
2
s
CueMCu 



The emf of the concentration cell will be



  1log
2
0591.002log
2
0591.00
MEME
cell
E

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MODULE-1
1
2
log
2
0591.0
M
M
cell
E 
The emf of the cell is + ve only if M2 > M1
The following characteristics of concentration cell can be noted:
1. When M2 = M1, the concentration cell does not generate electrical energy.
2. When M2 > M1, the Ecell is + ve.
3. When M2<M1, Ecell is – ve.
4. Higher the ratio of M2/M1, greater is the cell potential.
Types of electrodes:
1. Metal-Metal ion electrode: An electrode of this type consists of a metal dipped in a solution containing its
ions. Ex- Zn/Zn2+
, Cu/Cu2+
etc
2. Metal-Metal salt ion electrode: These electrodes consist of a metal is in contact with a sparingly soluble salt
of the same metal dipped in a solution containing anion of the salt.
Example-Calomel (Hg|Hg2Cl2|Cl-
, Silver- Silver salt electrode (Ag| AgCl |Cl-
)
3. Gas electrode: Gas electrode consists of a gas bubbling about an inert metal wire, immersed in solution
containing ions to which the gas is reversible. The metal provides electrical contact and facilitates the
establishment of equilibrium between the gas and its ions.
Example-Hydrogen electrode (Pt|H2|H+)
, Chlorine electrode (Pt|Cl2|Cl-
)
4. Oxidation-Reduction electrode: An oxidation-reduction electrode is a one in which the electrode potential
arises from the presence of oxidized and reduced forms of the same substance in solution. The potential arises
from the tendency of one form changes into the other more stable form. The potential developed is picked up by
an inert electrode like platinum.
Example-Pt|Fe2+
, Fe3+
Pt|Ce3+
, Ce4+
5. Ion selective electrode: In ion selective electrode, a membrane is in contact with a solution, with which it can
exchange ions.
Example-Glass electrode.
Reference Electrodes:
“Reference electrode are the electrode with reference to those, the electrode potential of any electrode
can be measured.”
It can acts both as an anode or cathode depending upon the nature of other electrode.
The Reference Electrodes can be classified in to two types
i) Primary reference electrodes Ex: Standard hydrogen electrode
ii) Secondary reference electrodes Ex: Calomel and Ag/Agcl electrodes

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MODULE-1
SHE has two main Limitations:
i) The construction of SHE is difficult. It is very difficult to maintain the concentration of H+
as 1M and
pressure H2 gas at 1atm
ii) Platinum electrode is poisoned by the impurities of the gas
Construction and working of Standard Calomel electrode (SCE):
1. Calomel electrode is a metal-metal salt Ion electrode.
2. It consists of mercury, mercurous Chloride and a solution of KCl. Mercury is placed at the bottom of a
glass tube.
3. A paste of mercury and mercurous chloride Is placed above the mercury. The space above the paste is filled
with a KCl solution of known concentration.
4. A platinum wire is kept immersed into the mercury to obtain electrical contact.
5. Calomel electrode can be represented as,
Hg | Hg2Cl2 | sat KCl
The calomel electrode can acts as anode or cathode depending on the nature of the other electrode of the cell.
The net cell reversible electrode reaction is,
Electrode potential,  20
log.
303.2 
 Cl
nF
RT
EE
 ,log.
303.20 
 Cl
F
RT
EE Where n=2
 Therefore electrode potential of calomel electrode is depending upon the concentration of KCl.
 The electrode is reversible with chloride ions.
 The potential of the calomel electrode depends on the concentration of the KCl.
 
 ClEE log0591.00
at 298K
Saturated KCl
Mercury
Calomel paste
Pt wire
Porous disc
Hg2Cl2(s) + 2e-
2Hg(l) + 2Cl-

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MODULE-1
For saturated KCl, the potential is 0.241V;
For 1M KCl , 0.280V;
For 0.1M KCl, 0.334V.
MEASUREMENT OF SINGLE ELECTRODE POTENTIAL USING CALOMEL ELECTRODE:
Electrode potential of a given electrode can be measured by using calomel electrode as a reference electrode.
Example-1: To measure the electrode potential of zinc: To measure the potential of the Zn- electrode, the
Zn- electrode is coupled with the SCE through a salt bridge. The anode and the cathode of the cell can be
identified by connecting the electrodes to the appropriate terminals of the voltmeter. Proper measurements can
be made only when the Zn-electrode is connected to the –ve terminal and the calomel electrode to the +ve
terminal of the voltmeter indicating that Zinc electrode is anode & the calomel electrode is a cathode.
HgClHgsatdKClSOZnZn 224 )(
At anode: Zn Zn2+
+ 2e
At cathode: Hg2Cl2 + 2e-
2Hg + 2C
Overall reaction: Zn + Hg2Cl2 Zn2+
+2Hg + 2Cl-
Ecell = Ecathode- Eanode
=ESCE –EZn
2+
/Zn
=0.2422V – Eo
Zn
2+
/Zn
EZn
2+
/Zn=0.2422-Ecell
EZn
2+
/Zn=0.2422-1.001
EZn
2+
/Zn= - 0.76V
Example-2: To measure the electrode potential of copper: Similarly to determine the copper electrode
potential of the cell, the cell is constructed as follows. Calomel electrode being the anode is connected to –ve
terminal of the voltmeter and copper electrode being the cathode is connected to the +ve terminal of the
voltmeter.
Hg/Hg2Cl2/KCl(sat)//Cu2+
/Cu
At anode: 2Hg + 2Cl-
Hg2Cl2 +2 e-
At cathode: Cu2+
+ 2e Cu
Overall reaction: 2Hg + 2Cl-
+ Cu2+
Hg2Cl2 + Cu
Ecell= Ecathode –Eanode
= Ecu
2+
/cu - ESCE
Ecu
2+
/cu = Ecell + 0.2422
Ecu
2+
/cu = 0.1 + 0.2422
Ecu
2+
/cu = +0.34V
Voltmeter
Calomel
Electrode
copper
Electrode
CuSO4
solution
Voltmeter
zinc
Electrode
Pt wire
ZnSO4
solutiom
Calomel
electrode

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MODULE-1
Advantages of calomel electrode:-
1. It is easily setup (simple to construct).
2. The cell potential is reproducible and stable over a long period.
3. It is used as a secondary reference electrode in the measurement of single electrode
potential.
4. It is the most commonly used reference electrode in all potentiometric determinations
and to measure pH of the given solution
Applications:
1. It is used as secondary reference electrode in the measurement of single electrode.
2. It is used as reference electrode in all potentiometer determinations and to measure pH of the given
solution.
Construction and working of Silver- Silver Chloride electrode:
1. Silver-Silver chloride is also a metal-metal salt ion electrode.
2. Silver and its sparingly soluble salt silver chlorides are in contact with a solution of chloride solution
ions. Generally a silver wire is coated with AgCl and dipped in a solution of KCl .
3. Cell representation is as follows
Ag |AgCl | sat KCl
Net half cell reaction is
AgCl + e-
Ag + Cl-
Electrodepotential  ,log.
303.20 
 Cl
nF
RT
EE Where n=1
 
 Cl
F
RT
EE log
303.20
 Therefore electrode potential of calomel electrode is depending upon
the concentration of KCl.
 The electrode is reversible with chloride ions.
 The potential of the calomel electrode depends on the concentration of the KCl.
For 1N solution, the electrode potential is 0.223V and for saturated solution is 0.199V at 298K
Applications:
1. Used as secondary reference electrode in ion selective elctrode.
2. In determining the distribution of potential on the ship hull and pipe lines.
 
 ClEE log0591.00
at 298K

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MODULE-1
Construction and working of Ion selective electrode (ISE):
“Ion selective electrode is one which selectively responds to a specific ion in a mixture and the potential
developed at the electrode is a function of the concentration of that ion in the solution”
Construction and working of Glass electrode:
A glass electrode is an ion selective electrode where potential depends upon the pH of the medium.
1. The glass electrode consists of a glass bulb made up of special type of glass (sodium silicate type of
glass) with high electrical conductance.
2. The glass bulb is filled with a solution of constant pH (0.1MHCl) and insert with a Ag-AgCl electrode,
which is the Internal reference electrode and also serves for the external electrical contact.
3. The electrode dipped in a solution containing H+
ions as shown in the figure.
4. The electrode representation is,
Glass | 0.1M HCl | Ag/AgCl.
INTERNAL SOLUTION EXTERNAL SOLUTION
C1= CONSTANT C2= [H+
]
E1 E2
Eb
The glass electrode works on the principle that when a thin glass membrane is in contact with a solution , A
boundary potential Eb is developed at layers of the glass membrane. This potential arises due to difference in the
concentration of H+
ion inside and outside the membrane.
Boundary potential, Eb = E2 – E1-------------------------(1)
1
Clog.
0591.00
1 n
EE 
2
log
0591.00
2
C
n
EE 
Where, C1 and C2 are concentration of H+
ions inner and outer membrane.
Substitute the values of E1 & E2 into eqn (1), we get





1
log
0591.00
2
log.
0591.00 C
n
EC
n
EEb
H+
Ion Solution
Ag/AgCl electrode
0.1 M HCl
GLASS ELECTRODE
GLASS
MEMBRANE

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MODULE-1
1
log
0591.0
2
log
0591.0
C
n
C
n
Eb  (n = 1)
1
log0591.0
2
log0591.0 CCEb 
(n=1, Since the concentration of the inner solution is constant, C1 is constant & (C2) = (H+
))



  HConstEb log0591.0
Where Const = K = -0.0591logC1
The glass electrode potential is sum of the
i) Boundary potential Eb,
ii) Ag-AgCl electrode potential EAg/AgCl and
iii) Asymmetry potential Easy.
asy
E
AgClAg
E
b
E
G
E 
/
Theoritically, Eb = 0 for C1 = C2. However, a small additional potential is exists called Easy
potential.
asy
E
AgClAg
EpHK
G
E 
/
0591.0
asy
E
AgClAg
EKGEwhere 
/
0
Determination of pH using glass electrode:
Procedure: glass electrode is immersed in the solution; the pH is to be determined. It is combined with a
reference electrode such as a calomel electrode through a salt bridge. The cell assembly is represented as,
pHKEb 0591.0
0591.00 pHGE
G
E 
pH Meter
Calomel
Electrode
Glass
Electrode
Solution of
Unknown pH

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MODULE-1
Hg| Hg2Cl2|Cl-
||Solution of unknown pH|glass|0.1M HCl|Ag|AgCl
The emf of the above cell, Ecell is measured using an electronic voltmeter with a pH meter.
The emf of the cell is given by
anode
E
cathode
E
cell
E 
……………… (1)
SCE
E
G
E
cell
E 
……………………… (2)
Since E SCE is knowing emf the cell,
E glass can be evaluated.
pHGE
G
E 0591.0
0
 …………………. (3)
SCE
EpHGEEcell  0591.0
0
…………..(4)
Advantages
1. This electrode can be used to determine PH
in the range 0-9, with special type of glass even up to 12 can be
calculated.
2. It can be used even in the case of strong oxidizing agents.
3. The equilibrium is reached quickly.
4. It is simple to operate, hence extensively used in various laboratories.
Limitations
1. The glass membrane though it is very thin, it offers high resistance. Therefore ordinary potentiometers cannot
be used; hence it is necessary to use electronic potentiometers.
2. This electrode cannot be used to determine the PH
above 12
1. Two Copper electrodes placed in CuSO4 solutions of equal concentration are connected to form o
concentration cell.
a) What is the cell voltage?
b) If one of the solutions is diluted until the concentration of Cu2+
ions is 1/5th
of its original value.
What will be the cell voltage after dilution?
Solution:
a) The cell potential of concentration cell is given as
1
2
log
0591.0
C
C
ncell
E 
When the concentration of the species are gqual (C2 = C1) the cell voltage is zero.
b) When one of the solution is diluted to
1
5
𝑡ℎ of its original value, C2 =1M & C1=
1
5
0591.0
0
SCEcellG EEE
pH



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MODULE-1
1
2
log
0591.0
C
C
ncell
E 
5/1
1
log
2
0591.0

cell
E
5log
2
0591.0

cell
E
699.002955.0 
cell
E
V
cell
E 0206.0
2. Two zinc rods are placed in 0.1M & 1M ZnSO4 solution separately to form a cell. Give the
electrochemical representation of the cell & calculate its emf.
Solution: Cell representation
Zn(S)/ZnSO4 (0.1M) ZnSO4 (1M)/Zn(S)
1
2
log
0591.0
C
C
ncell
E 
1.0
1
log
2
0591.0

cell
E
10log
2
0591.0

cell
E
V
cell
E 0295.0
3. Calculate the emf of the given concentration cell at 298K. Ag(s) /AgNO 3 (0.018M) AgNO3 (1.2M)
/Ag.
Solution :
1
2
log
0591.0
C
C
ncell
E 
018.0
2.1
log
1
0591.0

cell
E
66.66log
1
0591.0

cell
E
V
cell
E 1708.0
4. EMF of the cell Ag/AgNO3 (C1)// AgNO3 (C2=0.2M)/Ag is 0.8V. Calculate C1 of the cell.
Ecell = 0.0591/n logC2/ C1
0.8 = 0.591/1 log (0.2 / C1)
C1 = 5.5 X 10-14
M

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MODULE-1
5. The spontaneous galvanic cell Tin/Tin –ion (0.024M)//Tin-ion (0.064M)/Tin develops an Emf of
0.0126V at 25O
C. Calculate the valency of Tin.
Ecell = 0.0591/n logC2/ C1
0.0126 = 0.0591/n log (0.064/0.024)
n = 1.998 = 2.
Questions:
1. What is single electrode potential? Derive the Nernst equation for single electrode potential.
2. What are concentration cells? Deduce the expression for the EMF of a copper concentration cell.
3. Explain the construction & working of CALOMEL electrode.
4. Explain the measurement of electrode potential by using standard calomel electrode
5. Explain the construction & working of Ag/AgCl electrode
6. What is an ion selective electrode?
7. Explain the construction & working of GLASS electrode
8. Explain how glass electrode can be used in the determination of a PH
of a solution.

BATTERIES & FUEL CELLS
-----------------_____________________________________________________________________________________________
MODULE-1
Battery:
It is a compact device consisting of two or more galvanic cells connected in series or parallel or
both. It stores chemical energy in the form of active materials and on demand converts it into
electrical energy through redox reactions.
Batteries are used in calculators, digital watches, pace makers, hearing aids, portable computers,
electronically controlled cameras, digital watches, stand by power supplies, emergency lighting and
electroplating, telecommunication, military & space applications.
PRINCIPAL COMPONENT OF A BATTRY:
1. An anode where oxidation
2. A cathode where reduction occurs
3. An electrolyte , which is ironically conducting
4. A separator to separate anode and cathode compartments.
CLASSIFICATION OF BATTERIES
Batteries are classified into three types as follows.
a) Primary
b) Secondary
c) Reserve batteries.
a) Primary Batteries: These are the batteries which serve as a source of energy only as long as the
active chemical species are present in the battery or in the cell. The cell reactions are irreversible.
These are designed for only single discharge and cannot be charged again.
Ex: Dry Cell, Zn – HgO cell, Zn-air cell etc.
b) Secondary Batteries: These batteries are chargeable and can be used again and again. The cell
reactions are reversible and are often called reversible batteries. During discharging the cell acts like
galvanic cell converting chemical energy into electrical energy. During charging the cell acts like
electrolytic cell by converting electric energy into chemical energy, hence these batteries are called as
storage battery.
Ex: Lead acid Battery, Ni-Cd battery etc.
c) Reserve Batteries: The key components of the batteries such as electrolyte etc., is separated
from the rest of the component of the battery. And the battery is stored for a longer time. The
electrolyte is filled before its usage.
Ex: Mg – water activated batteries, Zn-Ag2O Batteries etc.
Operation of a battery during discharging and charging:
Discharge: During discharge, oxidation takes place at the anode and reduction takes place at the
cathode. The reaction is a spontaneous reaction. Chemical energy is converted into electrical energy.
The reactions occurring during discharge are
tcompartmencathodein
speciesactive
cathodenYMneYM
anodeneMM
c
nn
c
n
aa




At anode: electrons are released to the external circuit.
At cathode: electrons from the external circuit are consumed.

-----------------_____________________________________________________________________________________________
MODULE-1
Charging: During charging, reverse reactions take place. The reverse reactions are non-spontaneous
reactions. The battery is connected to an external d.c. power supply. Electrical energy is converted in
to chemical energy.
Example: The reverse of the above reactions occur during charging.
CHARACTERISTICS OF A BATTERY:
1. Voltage: The voltage of a battery mainly depends upon the emf of the cells which constitute the
battery system. The emf of the cell depends on the free energy changes in the overall cell reaction.
As given by Nernst equation,
 
 
 
 reactant
product
QquotientreactiontheisQandEEEwhere
Qlog
nF
RT303.2
EE
M
M
log
nF
RT303.2
EE
nFEG
anode
0
cathode
0
cell
0
0
n
0





Where Ecell =Ecathode- Eanode, and Q is the reaction quotient for the cell reaction at any stage of the
reaction. As it is evident from the above equation, is dependent on
a) Higher the standard electrode potential difference between the cathode and anode, higher is the
emf of the cell and the voltage available from the battery
b) As the temperature increases, emf of the cell decreases.
c) Emf of the cell decreases as the Q increases
2. Current: Current is a measure of the rate at which the battery is discharging. Higher the rate of
spontaneous reaction, higher is the current. Higher the surface area of the electrodes, higher is the
rate of reaction. Current is measured in A.
3. Capacity: Capacity is a measure of the amount of electricity that could be obtained from the fully
charged battery. It is expressed in Ah (ampere hours). It is proportional to the amount of charge in
Coulombs that may be transported from anode to cathode through the external circuit. The charge
(C) in Coulombs is given by the Faraday’s relation:
M
Fnw
C


Where, C is Capacity of battery (in Ah)
W is Weight of the active material
n is number of electrons involved in discharge reaction
F is Faradays constant, 96500 C/mol
M is Molar mass.

-----------------_____________________________________________________________________________________________
MODULE-1
4. Electricity storage density: It is the amount of electricity stored in the battery per unit weight of
the battery. i.e. it is the capacity per unit weight. It can be expressed in Coulombs/kg or in Ah/kg.
The weight includes the weight of all components of the battery (i.e. total weight of active
material, electrolyte, terminals etc.)
5. Energy efficiency: The energy efficiency of a rechargeable battery is given by
chargingduringconsumedEnergy
100gdischarginduringreleasedEnergy
efficiencyEnergy%


It holds good only for secondary battery.
6. Cycle Life: Primary batteries are designed for single discharge and secondary batteries can be
chargeable again and again.
The number of charge and discharge cycles that are possible in secondary batteries, before
failure occurs is called cycle life.
The cycle life of batteries must be high for secondary batteries.
7. Shelf life: The duration of storage battery under specified conditions at the end of which a cell or
a battery retains its ability performance lelvel is called shelf life. A good battery should have more
shelf life.
 Zn-air cell (Primary battery, non rechargeable)
Zinc –air battery is a modern and metal air battery. It uses oxygen from the atmosphere and it does not
contribute to the weight of the battery so these batteries offer high energy density.
Construction:
In zinc-air cell, granulated powder of zinc mixed with the electrolyte (KOH) acts as anode material.
Cathode is a carbon-catalyst mixture. The anode can and cathode can act as terminals. The anode
material is separated from the cathode material by an electrolyte absorbent separator. 5M KOH is used
as the electrolyte.
Anode: granulated zinc powder
Cathode: carbon – MnO2catalyst mixture
Electrolyte: 5M KOH
Separator: Polypropylene.

-----------------_____________________________________________________________________________________________
MODULE-1
Cell gives a voltage of 1.4V, Energy density 100Wh/Kg.
Cell representation:
air,CKOHZn
When air passed through the cell, zinc is oxidized to ZnO at the anode, during discharge.
Cell reactions:
At anode : Zn + 2OH-
At cathode : 1/2 O2 + H2O +2e-
Over all reaction Zn + 1/2 O2
ZnO + H2O + 2e-
2OH-
ZnO
Uses: Used in Military & radio receivers
Used as a power source in hearing aids.
Used in electronic pagers & various medical devices such as nerve & muscle simulator.
Used in drug impulsion equipment.
 Li-MnO2 cell: (primary battery, non rechargeable)
Li-MnO2 is a primary battery and produces a voltage of about 3V, Energy density 230Wh/Kg.
Lithium has the following advantages:
1. It is light.
2. It has a good electrical conductivity.
3. It has low standard electrode potential (Eo
= -3.05V) .
Construction:
1. The anode is made of lithium metal.
2. The cathode is made of MnO2.
3. A solution of lithium halide in organic solvent acts as the electrolyte.
4. The anode and cathode are separated by a polypropylene separator.
[Lithium halides: LiCl, LiBr, LiAlCl4
Organic solvents: Propylene carbonate and 1, 2 –Dimethoxyethane]
Cell representation : 2MnOsolventsorganicinhalideLithiumLi

-----------------_____________________________________________________________________________________________
MODULE-1
Cell reactions:
At anode : Li
At cathode : MnO2 + Li+
+ e-
Over all reaction: Li + MnO2
Li+
+ e-
LiMnO2
LiMnO2
Mn (IV) reduced to Mn( III) & Li+
enters crystal lattice of MnO2
Uses:
Used as memory back up equipments.
Used in watches, calculators, toys, cameras etc.
Used in safety & security devices
 LITHIUM-ION BATTERIES
The lithium ion batteries are rechargeable battery best suited to mobile devices that requires small size,
light weight and high performance.
In lithium-ion batteries, lithium compounds are used as anode. These batteries are known as re-
chargeable batteries. Therefore, Lithium ion batteries are considered as best than pure Lithium based
batteries. It works on the principal of Intercalation mechanism.
CONSTRUCTION:
1. Li-ion cell has a four-layer structure.
2. Anode: Lithium intercalated graphite/carbon (specialty carbon)
3. Cathode : lithium metal oxide compound such as LiyNiO2 , LiyCoO2 and LiyMnO2
4. Anode current collector -copper foil
5. Cathode current collector- aluminum foil
6. Separator : Polypropylene
7. An electrolyte made with lithium salt [LiPF6] in an organic solvent [propylene carbonate or 1,2 –
dimethoxyethane]..
8. Lithium ion secondary battery depends on an "intercalation" mechanism.

-----------------_____________________________________________________________________________________________
MODULE-1
Cell reactions :
During discharge Li ions are dissociated from the anode and migrate across the electrolyte and are
inserted into the crystal structure of the host compound of cathode.
During charging, lithium in positive electrode material is ionized and moves from layer to layer and
inserted into the negative electrode.
At the same time the compensating electrons travel in the external circuit and are accepted by the host
to balance the reaction.
Advantages
1. They have high energy density than other rechargeable batteries
2. They are less weight
3. They produce high voltage out about 4 V as compared with other batteries.
4. They have improved safety, i.e. more resistance to overcharge
5. No liquid electrolyte means they are immune from leaking.
6. Fast charge and discharge rate
Applications:
1. The Li-ion batteries are used in cameras, calculators.
2. They are used in cardiac pacemakers and other implantable device
3. They are used in telecommunication equipment, instruments, portable radios and TVs, pagers
4. They are used to operate laptop computers and mobile phones and aerospace application.
 Nickel-metal hydride battery
[Alkaline storage battery & Secondary battery]
Construction:
1. In these batteries, electrodes are made of porous nickel foil or nickel grid, into which the active
material is packed.
2. Anode: The active material for the anode is a mixture of a metal hydride (such as TiH2, VH2, or
ZrH2) with a hydrogen storage alloy ( such as LaNi5 or TiNi ).
3. Cathode: The active material for cathode is nickel oxy hydroxide, NiO(OH).
4. Electrolyte: An aqueous solution of KOH acts as the electrolyte.
5. Separator: Polypropylene
Cell representation : 2i(OH)N/NiO(OH)/MH/KOH(5M)
The battery produces 1.25 to 1.35V per cell.

-----------------_____________________________________________________________________________________________
MODULE-1
Cell reactions:
At anode : MH + OH-
At cathode : NiO(OH) + H2O + e-
Over all reaction: MH + NiO(OH)
M + H2O + e-
Ni(OH)2 + OH-
M + Ni(OH)2
Uses: Used in cellular phones, camcorders and laptop computers.
Fuel cells:
Fuel cells is defined as a
Galvanic cells in which chemical energy of a fuel directly converted into electrical energy.
Basic component of fuel cell:
1. Fuel cells consist of electrodes and electrolytes.
2. Catalyst used is embedded in the electrodes.
3. Gaskets are used to prevent the leakage of gases between the electrodes.
Reactions:
At anode: Fuel  Oxidised product + ne-
At cathode: Oxidant + ne-
 Reduced product
Fuel cells are represented as: Fuel /anode/electrolyte/cathode/oxidant
Difference between a battery and a fuel cell:
BATTERIES FUEL CELLS
Batteries are energy storage devices Fuel cells are energy conversion devices.
Secondary batteries are rechargeable Fuel cells are not chargeble.
The reactants and products form an integral
part of batteries.
In fuel cells, continous movement of fuel,
oxidant, and reaction product in and out of cells.
Advantages:
1. High power efficiency approximately 75%.
2. No need of charging.
Cell can
Cathode
Separator
Anode
Sealing washer
separator
Cell cap
Contact spring

-----------------_____________________________________________________________________________________________
MODULE-1
3. Produces direct current for a long time.
4. No moving parts. Hence wear and tear is eliminated.
5. Harmful products are absent. Hence fuel cells are ecofriendly.
Limitations:
1. Electrodes and electrolytes are expensive.
2. Storage of fuel and oxidant.
3. Gives DC output and should be converted to AC.
Methanol – Oxygen fuel cell
It is an electroconductive organic fuel cell .
Construction:
1. It consists of anode and cathode made of platinum.
2. Sulphuric acid acts as the electrolyte.
3. A membrane is inserted adjacent to the cathode on the electrolyte side to minimize the diffusion
of methanol into the cathode.
4. Methanol – H2SO4 mixture is circulated through the anode chamber.
5. Pure oxygen is passed through the cathode chamber.
Cell reactions:
At anode : CH3OH + H2O
At cathode : 11/2 O2 + 6H+
+ 6e-
Over all reaction: CH3OH + 11/2 O2
3H2O
CO2 + 2H2O
CO2 + 6H+
+ 6e-
Uses: It is used in large-scale power production.
Used in space vehicles, military & mobile power systems.
O2CH3OH +
H2SO4
Excess Oxygen and
water
Cathode
Membrane
H2SO4 (electrolyte)
Anode
CO2
CO2

-----------------_____________________________________________________________________________________________
MODULE-1
Questions:
1. What is battery? Explain the classification of battery.
2. Explain the following battery characteristics:
i. voltage, ii. Current , iii. Capacity, iv. Electricity storage density, v. Energy efficiency ,
vi. Cycle life vii. Shelf life
3. Explain the construction, working & applications of ZINC-AIR battery
4. Explain the construction, working & applications of Li-MnO2 cell
5. Explain the construction, working & applications of Li-Ion battery
6. Explain the construction, working & applications of Ni-MH battery
7. What is a fuel cell? How is it different from the conventional batteries
8. Explain the construction, working & applications of METHANOL-OXYGEN fuel cell

Module 1 electrode potential &amp; cells - battery technology

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