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Electrochemistry: It is a branch of chemistry that studies chemical reactions which take place in a solution at
the interface of an electron conductor (a metal or a semiconductor) and an ionic conductor (the electrolyte), and
which involve electron transfer between the electrode and the electrolyte.
Electrolyte: it is a substance which is an aqueous solution or molten state liberates ions and allows current to
flow through it.
Electrolysis: when electric current passes through the electrolyte which results in chemical decomposition and
this phenomenon is called electrolysis.
Electrolytic cells: The conversion of electrical energy to chemical energy takes place. Electrolysis takes place
here. It contains aqueous solution of an electrolyte in which two metallic rods are dipped which is connected to a
battery.
Anode: The electrode through which current enters the cell is known as anode. It is denoted as the positive
electrode.
Cathode: The electrode through which current leaves the cell. It is denoted as the negative electrode.
Ohm’s Law: The current “I” flowing through a conductor is given by relation E/R where E is the electromotive
force and R is the resistance.
I =E/R
(OR)
Ohm's law states that the current passing through a conductor between two points is directly proportional to the
potential difference or voltage across the two points, and inversely proportional to the resistance between them.
The mathematical equation that describes this relationship is
Where I is the current, Ampere in units
Electromotive Force: The potential which is required to move a unit charge from one place to another place. It
is also known as E.M.F It is measured in volts.
Volt: The volt is defined as the value of the voltage across a conductor when a current of one ampere strength
through a one ohm resistance.
Amperes: The unit of strength of current is known as amperes. The current which deposits 0.001118 gms of
silver per second from a 15% solution of silver nitrate in Voltameter.
Ampere = volt/ohm
Coloumb: The quantity of current is measured in Coloumb.
The quantity of current which passes in one second with a current strength of one ampere.
Electrical resistance: The electrical resistance of an object is a measure of its opposition to the passage of an
electric current. The unit of electrical resistance is the ohm (Ω). Resistance’s reciprocal quantity is electric
conductance is measured in Siemens.
The resistance of an object can be defined as the ratio of voltage to current:




Three types of conductance
Specific Conductance
Equivalent Conductance
Molecular conductance
 Specific Conductance: The resistance offered by a conductor to the passage of electricity through it is
directly proportional to length and inversely proportional to the area of cross section. The resistance R is
given by the relation:
A certain weight of an electrolyte is dissolved in Vml of solvent and the conductance of one ml of the resulting
electrolyte solution at a given dilution V is called the Specific Conductivity.
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 Equivalent conductance:
If one gram equivalent weight of an electrolyte is dissolved in Vml of the solvent, the conductivity of all ions
produced from one gram equivalent of an electrolyte at the dilution V is known as Equivalent Conductance. It is
denoted by λv. Hence the equivalent conductance is equal to the product of specific conductance and volume.
λv = Kv x v = Kv x 1000/ N
Units are ohm-1 cm2 equiv-1
 Molar Conductance:
The conductance of all ions produced by dissolving one gram molecular weight of one mole of an electrolyte
when dissolved in a certain volume Vml. Molar conductance is denoted by µv.
µv= Kv x V =
Kv x 1000/ M
Units are ohm-1 cm2 mole-1
 EFFECTS OF DILUTION ON CONDUCTANCE:
Due to dilution ionization increase and specific conductance decreases. This happens because specific
conductance is the conductance of the ions present in one centimeter cube of the solution. On dilution the
number of current carrying particles (or) ions presents per one centimeter cube of the solution decreases.
Equivalent conductance or molecular conductance of an electrolyte increases on dilution because
 The above conductance is the product of kv and the volume of the solution. When volume increases
equivalent conductance also increases.
 The number of ions of the electrolyte solution increases on dilution contributing to the increase of
conductance. Ionization increases on dilution, till the whole electrolyte substance has ionised. the limiting
value is known as equivalent conductance at infinite dilution and it is represented by the symbol λα.
The conductance ratio is called the degree of ionization:
α = λv /λα
 Electrolytes can be divided into two types:
 Strong Electrolytes
 Weak electrolytes
 Strong Electrolyte:
A strong electrolyte is a substance that gives a solution in which almost all the molecules are ionized, even at
low concentration such solution has increasing value of equivalent conductance at low dilution.
Strong Acids: HCl, H2SO4, HNO3
Strong Base: NaOH, KOH
The Salts: Practically all salts are strong electrolytes
 Weak Electrolyte:
The electrolyte which ionize to a small extent on dilution are called weak electrolyte. They have a low value of
equivalent conductance even at a higher concentration and are not completely ionized even at very great dilution.
Weak Acids: All organics acids like acetic acids, propionic acid and H2SO3
Weak Bases: Alkyl Amines, NH4OH
Salts: A few salts such as mercuric chloride and lead acetate
 Measurement of conductance :
Conductance is the reciprocal of resistance and the resistance can be determined by a Wheatstone bridge circuit
in which the conductivity cell forms one arm of the bridge. Wheatstone bridge circuit for is used for measuring
conductivity, Conductivity cell with one arm of a resistance bridge for measurement of conductivity of an
electrolyte.
The conductance cell is made of highly resistant glass such as Pyrex or quartz. The electrodes consist of
platinum discs coated with finely divided platinum black. These are called Platinised platinum electrodes.
Platinum black surface catalyze the union of hydrogen and oxygen which tend to be liberated by the successive
pulse of the current and the polarization E.M.F is thus eliminated. The electrodes are welded to platinum wires
fused in two glass tubes. The glass tube contains mercury and is firmly fixed in the ebonite cover of the cell, so
that the distance between the electrodes may not alter during the experiment.
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The cell is connected to a Wheatstone bridge which consists of a wire of platinoid or manganin AC having a
uniform thickness so that the ratio of lengths read on the scales gives the ratio of resistance. The wire of AC is
stretched tightly over a meter scale graduated in millimeters. A sliding contact D moves along the wire. R is the
resistance box and B is the conductance cell. Conductance cell containing electrolyte is placed in the thermostat
for maintaining constant temperature during the measurement of conductance. The induction coil is used to pass
alternate current in the circuit. The sliding contact D is moved until the sound in the head phones is minimum.
This gives the null point where resistance of the resistance box R and resistance of the electrolyte in the cell B
are equal.
The arms AD and DC represented by resistance R1 and R2 are usually in the form of a single calibrated slide
wire resistor with a sliding contact connected to the null detector. The solution whose conductance is to be
determined is placed in conductivity cell. When the bridge is balanced, assuming that the conductivity cell
behaves as a pure resistance, then the voltages between AD from resistance box and DC from conductance cell
are equal i.e. No sound is observed in the ear phone.
Mathematically it is expressed as “at null point”
Resistance of resistance box = resistance of AD
(1)
Resistance of solution in conductance cell = resistance DC
(2)
Divide the eq (1) by (2)
Resistance of resistance box (R)
= resistance of AD
Resistance of solution in conductance cell(C)
resistance DC
= length of AD
length of DC
1
=
length of AD × 1
Resistance of cell (C)
length of DC
R
Thus Observed conductance = length of AD × 1
[since observed conductance = 1/ Resistance of cell (C)]
length of DC R
 Determination of Cell Constant:
The electrodes in the cell are not exactly 1 cm apart and may not have surface area of 1 sq. cm (1cm2). Thus the
value of observed conductivity is not equal to specific conductance but is proportional to it.
R=ρ l/A, C=1/R = (A/l)×(1/ ρ)
Let l/A = x = cell constant,
let 1/R = observed conductance
We know that
1/ ρ = Specific conductance
Thus Cell constant (x) = specific conductance / observed conductance
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 KOHLARAUSH LAW:
The equivalent conductance at infinite dilution of different electrolytes is the sum of the ionic conductances of
cations and anions.
λα = λa + λc
Ionic conductance is expressed in ohm-1 cm2 equiv-1
Ionic conductance is directly proportional to the transport numbers.
λa = K x v (1)
λc = K x u
(2)
λα = λa + λc =K (u +v)
(3)
Divide the (1)/(3) & (2)/(3), we get
Let v/(u+v) = n and u/ (u+v) =1-n
λa / λα = v/(u+v) = n
λc / λα = u/ (u+v) =1-n
 Applications of Kohlrausch’s Law:
a) Calculation of conductivity of a weak electrolyte at infinite dilution
It is not possible to determine the value of λα for weak electrolytes since we cannot obtain the limiting value of
the molar conductivity for a weak electrolyte. This is done indirectly by the molar ionic conductance for the
individual ions of the weak electrolyte and by using Kohlrausch’s law.
λa/ λc = n/1-n => λa - n λa =n λc
=> λa = n (λa+ λc ) => λa = n λα
Similarly
λc = (1-n) λα
b) Determination of Solubility of Sparingly Soluble Salts
Salts like AgCl, BaSO4, etc. are sparingly soluble salts and have a very small but definite solubility in water.
The solubility of such sparingly soluble salts is obtained by determining the specific conductivity (κ) of a
saturated salt solution.
λv = Kv x V = λα (= λa + λc)
c) Calculation of Apparent Degree of ionization or conductivity ratio
The apparent degree of ionization or conductivity ratio of an electrolyte (α) is the ratio of equivalent conductance
of a solution at definite dilution to the equivalent conductance of a solution at infinite dilution.
i.e.,
α = λv/λα
 Conductometric titrations:
Conductometric means measuring the conductivity of ionic solutions caused by mobility of ions towards
respective electrodes in the presence of electric field.
Definition: The determination of end point of a titration by means of conductivity measurements are
known as Conductometric titrations.
 Types of Conductometric titrations:
a) Titration of strong acid with strong base (e.g. HCl with NaOH)
Consider a reaction in which HCl is titrated against NaOH. Pippet out 20ml of acid solution is taken
in a beaker in which a conductivity cell is dipped. Add 1ml of NaOH to the beaker solution (HCl)
and measure the conductance by using conductometer to which conductivity cell is attached.The
process is repeated after every addition of NaOH, plot a graph between Volume of base (X-axis) and
Conductance (Y-axis) we get a V-shaped curve which denotes that the conductance of HCL is
slowly decreased by the addition of NaOH up to a stage and then suddenly increase the conductance
after end point.
HCl + NaOH  NaCl + H2O
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b) Titration of weakacid with strong base (e.g. CH3COOH with NaOH)
CH3COOH + NaOH
 CH3COONa + H2O
When a weak acid (CH3COOH) is titrated against a strong Base (NaOH), it gives low conductance due to poor
dissociation of the weak acid. On adding alkali, highly ionized sodium acetate is formed. Later on the
conductance is increased due to highly ionized salt (is the end point). After it the addition of NaOH contributes
sharp increase in the conductivity.
c) Titration of strong acid with weak base base (e.g. HCl with NH4OH)
In this case the conductivity of the solution will first decrease due to fast moving H+ ions by slow
moving NH4+ ions. After end point conductivity has no change in it.
d) Titration of weakacid with strong base (e.g. CH3COOH with NH4OH)
CH3COOH + NH4OH
 CH3COONH4 + H2O
In this case the conductivity of the solution will first decrease due to poor dissociation of the weak
acid, but starts raising as CH3COONH4 is formed. After end point the conductivity remains almost
constant, because the free base NH4OH is a weak electrolyte.
 Galvanic cells: A device which converts chemical energy to electrical energy is called Electrochemical
cell (or Galvanic cell or voltaic cell)
EMF of such cell is directly proportional to intensity of the chemical reaction taking place in it.
A typical example for Galvanic cell is DANIEL Cell.
DANIELL CELL
It consists of copper rod dipping in CuSO4 solution, Zinc rod is dipping in ZnSO4 and two solutions are
connected by metallic wire and separated by a semi permeable membrane or salt bridge. It only allows the
passage of ions but not solutions.
When two electrodes are connected by wire flow of current takes place by the high oxidation potential of Zn rod,
it releases the electrons passes through the wire and forms metallic copper at Cathode.
The cell reactions are
Zn(s)
 Zn2+ + 2e at anode
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Cu2+ + 2e  Cu
at cathode
Over all reaction
Zn(s) + Cu2+

Cu(s) + Zn2+(aq)
The cell is represented as
Zn / ZnSO4 // CuSO4 / Cu , emf of cell= 1.09 Volts
The negative electrode is extreme left and the positive electrode is extreme right. The double vertical line
between two liquids represents salt bridge.
The process of a metal passing metal ions in to solution by liberating electrons is called de-electro nation or
oxidation or Anode.
Zn(s)
 Zn2+ + 2e at anode
The process of gaining electrons is called electro nation or reduction or Cathode.
Cu2+ + 2e  Cu
at cathode
The anode is negative sign because the current leaves from the cell and cathode is positive sign because
current enters into the cell.
The emf of the Daniel cell is 1.09 volts.
Note:
a) If we apply external emf which is greater than 1.09 volts , then the cell reactions are reversed like
Cu + Zn2+  Cu2+ + Zn
b) If we apply an emf of less than 1.09 volts, then no change in the cell reaction takes place.
c) If we apply an emf of exactly 1.09 volts then the reactions are stops.
Any cell which does not satisfy the conditions is Irreversible, otherwise called as Reversible cell.
 Half cell:
An Electrode which is dipped in its salt solution is called half cell or single electrode cell and the potential which
is developed between the electrode and its salt solution is called single electrode potential.
The standard hydrogen half cell:
2H+(aq) + 2e- → H2(g)
The half cells of a Daniel cell:
Half cell (anode) of Zn
Half cell (cathode) of Cu
Zn + Cu2+
Zn
Cu2+ + 2e−
→ Zn2+ + Cu
→ Zn2+ + 2e−
→ Cu
Measurement of single electrode potential
i) Emf of the cell = EAnode + ECathode
= Oxidation potential of anode (EL) + Reduction potential of cathode (ER)
ii) With respect to reduction potentials Emf of the cell (Ecell) = ER - EL
iii) If the standard electrode is acts as Anode then
iv) If the standard electrode is acts as Cathode then
Ecell = ER – E0
Ecell = E0 – EL

DERIVATION OF NERNST EQUATION: The electrode potential and the emf of the cell depend
upon the nature of the electrode, temperature and the activities (concentrations) of the ions in solution. The
variation of electrode and cell potentials with concentration of ions in solution can be obtained from
thermodynamic considerations. For a general reaction such as the Gibbs free energy change is given by the
equation
G = ∆Go + 2.303RT log10 [P/R]....... (i)
o
where ∆G refers to free energy change for the reaction when the various reactants and products are present at
standard conditions. The free energy change of a cell reaction is related to the electrical work that can be
obtained from the cell, i.e.,
∆G = -nFEcell
and
∆Go = -nFEo.
On substituting these values in Eq. (i) we get
-nFEcell = -nFEo + 2.30eRT log10 [P/R]....... (ii)
or
Ecell = Eocell - 2.303RT/nF log10 [P/R]....... (iii)
Substituting the values of R=8.314 JK-1 mol-1, T = 298 K and F=96500 C, Eq. (iv) reduces to
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E = Eo - 0.0591/n log10 ....... (v)
= Eo - 0.0591/n log10 ([Products])/([Reactants]) ....... (vi)
This equation is known as Nernst equation.
 Type of electrodes(Reference electrodes): These are also called as Standard electrodes and used
for the determination of EMF of a cell. The following are the types
a) standard hydrogen electrode (SHE)
It is also called as Normal Hydrogen Electrode. SHE Consists of a Platinum Electrode coated with
platinum black immersed in a 1M solution of H + ions maintained at 25 0 c. Hydrogen gas of 1 atm
pressure is enters the glass hood and when it contact with the platinum foil which is attached with
platinum Electrode release the electrons. These electrons consumed by the H + ions present in the
solution to which the foil is in contact and Hydrogen gas liberated as Bubbles from the solution.
The cell reactions are
½H2
 H+ + e
(at foil or anode)
+
H + e

½H2
(at solution or cathode)
By the reactions we can say that the amount of gas is inserted into the glass hood is liberated as gas from the
solution i.e. there is no chemical (or potential) change occur. So the potential developed in this cell is Zero.
When another electrode cell which has a solution of unknown concentration is attached with SHE, the emf
developed in the arrangement is calculated
Ecell = ER - EL where ER= right side electrode EL= left side electrode
Ecell = Eocell - 2.303RT/nF log10 [P/R]
= 0 – 0.0591 log(H+) (SHE emf is zero)
Ecell = 0.0591 PH
By using this equation we can calculate the concentration of any solution which is taken in the cell.
b) Standard Calomel Electrode
Calomel electrode: It consists of a glass tube, having a side tube on each side. The mercury is placed at the
bottom over which a paste of mercury-Mercurous chloride is placed. A solution of potassium chloride is then
placed over the paste. A platinum wire sealed in a glass tube helps in making the electrical contact. The
electrode is connected with the help of the side tube on the left through a salt bridge with the other electrode to
make a complete cell and measure the potential.
The potential of the calomel electrode depends upon the concentration of the potassium chloride solution. The
cell is represented as Hg, Hg2Cl2/KCl.
For saturated KCl solution electrode potential is 0.2415 volts
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For 1N KCl solution electrode potential is 0.28 volts
For 0.1 N KCl solution electrode potential is 0.3338 volts
If we connect an electrode which is dipped in an solution of unknown concentration with
standard calomel electrode, then the emf developed in this cell is calculated as
Ecell = ER - EL
ER= right side electrode (Calomel)
EL= left side electrode (unknown)
Ecell = Eocal - 2.303RT/nF log10 [P/R]
Ecell
= 0.2415 – 0.0591 log (H+)
Ecell = 0.2415 + 0.0591 PH (for saturated KCl solution)
Ecell = 0.28 + 0.0591 PH (for 1N KCl solution)
Ecell = 0.3338 + 0.0591 PH (for 0.1 N KCl solution)
c) ION- Selective Electrode(ISE): ISE use a membrane which is sensitive to a particular chemical species
i.e. it respond to specific ions in a solution while leaving others in solution to develop potential. In ISE
the use of membrane is must. Generally following membranes are used
i) - Glass electrodes
ii) - Liquid electrodes
iii) - Solid electrodes
1. Solid-state electrodes based on inorganic salt crystals used in the form of pressed pallets of Ag2S,
AgCl in combination with lanthanum fluoride.
2. Liquid-based electrodes using a hydrophobic polymer (10-phenonthralene Fe solution) membrane
saturated with a hydrophobic liquid ion exchange.
The membrane potential is calculated as E= (RT/nF) Log (C2/C1)
Suppose the reference electrode is anode and the membrane is cathode then
Ecell =E - Eref here E ref is zero if identical eletrodes were used as the concentration of ions in solutionin
which reference electrode is dipped(C2) is Constant.
3. Glass electrode:
The glass electrode used to measure pH is the most common ion-selective electrode. The glass electrode is made
of Na2O, CaO and SiO2.
A typical pH combination electrode, incorporating both glass and reference electrodes (Calomel) in one body.
Glass combination electrode with a silver-silver chloride reference electrode.
When we insert the electrode in silver solution, the emf developed is calculated as EG = EG0 + 0.0591 PH
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The glass electrode is immersed in a solution of unknown pH so that the porous plug on the lower right is below
the surface of the liquid. The two silver electrodes measure the voltage across the glass membrane. A glass
electrode is a type of ion selective electrode consists of a thin walled glass bulb attached to a glass tube. A very
low melting point and high electrical conductivity glass is used for the construction of this bulb. The glass tube
contains a dilute solution of constant pH of HCl (0.1 N) solution. A silver-silver chloride electrode or platinum
wire is immersed as reference electrode in the HCl solution. The working of glass electrode is based upon the
observation that when a glass surface is in contact with a solution, there exists a potential differencebetween the
glass surface and the solution, the magnitude of which depends upon the H+ ion concentration of the solution and
the nature of glass. The glass electrode may be represented as
Ag, AgCl(s) | 0.1 N HCl | glass | H+ = unknown
The electrode potential of the glass electrode depends upon the concentration of H+ ions contained in the
experimental solution and are given
The emf is calculated as
Ecell = ER - EL
ER= right side electrode (Calomel)
EL= left side electrode (unknown with glass)
Ecell = Eocal - EG
Ecell = Eocal – (EG0 + 0.0591 PH)
Ecell
= 0.2415 – EG0 - 0.0591 log (H+)
0
Ecell = 0.2415 – EG - 0.0591PH (for saturated KCl solution)
Ecell = 0.28 – EG0 - 0.0591PH (for 1N KCl solution)
Ecell = 0.3338 – EG0 - 0.0591PH (for 0.1 N KCl solution)

Concentration Cells:
If two electrodes of same metal are dipped separately in two solutions of same electrolyte and are connected
with a salt bridge, the whole arrangement is found to act as a galvanic cell.” A concentration cell is a Galvanic
cell in which electrical energy is produced by the transfer of material from a system of higher concentration to one
of low concentration”. In general, there are two types of concentration cells
(i) Electrode concentration cells:
In these cells, the potential difference is developed between two like electrodes at different concentrations dipped
in the same solution of the electrolyte. For example, two hydrogen electrodes at different pressure in the same
solution of hydrogen ions constitute a cell of this type.
(Pt,H2 (Pressure p1))/Anode |H+ | (H2 (Pressure p2)Pt)/Cathode
The cell reactions are H2 (P1)
 2H+ + 2e (oxidation)
2H+ + 2e  H2(P2)
(Reduction)
Total reaction
H2 (P1)
 H2(P2)
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If p1, p2 oxidation occurs at L.H.S. electrode and reduction occurs at R.H.S. electrode. i.e. it is clear that in this
process there is no net chemical change is occurred but only change in concentration takes place at two
electrodes.
Ecell = (0.0591/2) log(P1/P2) at 25o C
In the amalgam cells, two amalgams of the same metal at two different concentrations are interested in the same
electrolyte solution. For example two unequal concentration of zinc amalgam dipped in a solution of ZnSO4.
The cell is represented as
Zn(Hg) C1  Zn(Hg) C2
The emf is calculated as
Ecell = (0.0591/2) log(C1/C2) at 25o C
(ii) Electrolyte concentration cells:
In these cells, electrodes (Zn) are identical but these are immersed in solutions of the same electrolyte (ZnSO4)
of different concentrations (C1 & C2). The source of electrical energy in the cell is the tendency of the electrolyte
to diffuse from a solution of higher concentration to that of lower concentration. With the expiry of time, the two
concentrations tend to become equal. Thus, at the start the emf of the cell is maximum and it gradually falls to
zero. Such a cell is represented in the following manner:
(C2 is greater than C1).
Zn|Zn2+(C1)||Zn2+(C2)|Zn
or
(Zn|Zn2+ (C1))/Anode || (Zn2+ (C2 )|Zn)/Cathode
The emf of the cell is given by the following expression:
Ecell = (0.0591/n) log C(2(R.H.S.))/C(1(L.H.S.)) at 25o C
The concentration cells are used to determine the solubility of sparingly soluble salts, valency of the cation of the
electrolyte and transition point of the two allotropic forms of a metal used as electrodes, etc.

BATTERIES
Definition:
 A battery is a storage device used for the storage of chemical energy and for the transformation of
chemical energy into electrical energy
 Battery consists of group of two or more electric cells connected together electrically in series.
Battery acts as a portable source of electrical energy.
Energy produced by an electrochemical cell is not suitable for commercial purposes since they use salt
bridge which produces internal resistance which results in drop in the voltage. The drop in voltage is negligible
only for a small interval of time during which it is being used.
Batteries are of 3 types. Namely
• Primary Batteries (or) Primary Cells
• Secondary Batteries (or) Secondary Cells
• Fuel Cells (or) Flow Batteries
I. Primary Batteries (or) Primary Cells:-
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Primary cells are those cells in which the chemical reaction occurs only once and the cell becomes dead after
sometime and it cannot be used again. These batteries are used as source of dc power.
Eg. Dry cell (Leclanche Cell) and Mercury cell, lithium cell.
Requirements of Primary cell:
It should satisfy these requirements
1) It must be convenient to use.
2) Cost of discharge should be low.
3) Stand-by power is desirable.
Dry cell (Leclanche Cell)
It consists of a cylindrical Zinc container that acts as an anode. A graphite rod placed in the centre (but not
touching the base) acts as a cathode. The space between anode and cathode is packed with the paste of NH4Cl
and ZnCl2 and the graphite rod is surrounded by powdered MnO2 and carbon as shown in Figure. The cell is
called dry cell because of the absence of any liquid phase, even the electrolyte consists of NH4Cl ,ZnCl2 and
MnO2 to which starch is added to make a thick paste which prevents leakage. The graphite rod is fitted with a
metal cap and the cylinder is sealed at the top with a pitch.
The Zn-MnO2 cell (dry cell) is represented as Zn/Zn+2,NH4+/MnO2/C (EMF = 1.5V)
At anode: (Oxidation)
Zn(s) 
Zn+2(aq)+2eAt Cathode: (Reduction)
2MnO2(s) +H2o+2e-  Mn2O3(s) +2OH –
The net cell reaction is
Zn(s) +2MnO2(s) +H2O Zn2++ Mn2O3+ 2OHThe resulting OH- ions react with NH4Cl to produce NH3 which is not liberated as gas but immediately
combines with the Zn2+ and the Cl- ions to form a complex salt [Zn(NH3)2Cl2] (diammine dichloro zinc).
2 NH4Cl + 2 OH-  2NH3 + Cl- +2 H2O
Zn2+ + 2NH3 + 2 Cl-  [Zn (NH3)2Cl2]
Advantages:
1) These cells have voltage ranging from 1.25v to 1.50v.
2) Primary cells are used in the torches, radios, transistors, hearing
aids, pacemakers, watch etc.
3) Price is low.
Disadvantages:
These cells do not have a long life, because the acidic NH4Cl corrodes the container even when the cell is not
in use.
Lithium cells:Lithium Cells are Primary cells in which lithium acts as anode and cathode may differ. Lithium metal is used as
anode because of its light weight, high standard oxidation potential (>3v) and good conductivity.
P a g e | 12
As the reactivity of lithium in aqueous solution is more, Lithium cells use non aqueous solvents as electrolyte.
Lithium cells are classified into two categories:
a) Lithium cells with solid cathodes
b) Lithium cells with liquid cathodes
(a) Lithium cells with solid cathode: The electrolyte in these systems is a solid electrolyte most widely used
cell is Lithium-Manganese dioxide cell(3V) MnO2 should be heated to over 3000C to remove water before
keeping it in the cathode, there by the efficiency of the cell is increased.
Anode: Lithium metal
Cathode: MnO2 as an active material
Electrolyte: LiBF4 salt in a solution of propylene carbonate and dimethoxy ethane
• Reactions:
At Anode:
Li  Li+ +eAt Cathode: e + MnO2  MnO2Net reaction:
Li+ + MnO2-  Li MnO2
Applications: 1) The coin type cells are used in watches and calculators
2) Cylindrical cells are used in fully automatic cameras.
(b) Lithium cells with Liquid cathode: Lithium- Sulphur dioxide cell is an example of liquid cathode. The cosolvents used are acrylonitrile or propylene carbonate (or) mixture of the two with SO2 in 50% by volume.
Cell reaction: 2Li + 2SO2 → LiS2O4
Lithium thionyl chloride cell is another example of liquid cathode. It consists of high surface area carbon
cathode, a non-woven glass separator. Thionyl chloride acts as electrolyte and as cathode.
Cell Reaction:
At Cathode:
4Li → 4Li+ + 4e+
At Anode:
4 Li + 4e- + 2 SOCl2 → 4 LiCl + SO2 +S
4 Li + 2 SOCl2 → 4 LiCl + SO2 +S
In this cell no co- solvent is required as SOCl2 is a liquid with moderate vapour pressure.
The discharging voltage is 3.3- 3.5 V.
Uses: 1) They are used for military and space application.
2) In Medicinal devices like neuro-stimulators drug delivery system lithium batteries are widely used.
3) They are also used in electric circuit boards for supplying fixed voltage for memory protection and standby
functions.
II. Secondary Cells (or) Accumulator batteries:These cells can be recharged by passing an electric current through them and can be used again and again.
Eg: A. Lead storage battery
B. Nickel-Cadmium battery
Secondary cells are widely used in cars, trains, motors, electric clocks, power stations, laboratories, emergency
lights, telephone exchange, digital cameras, laptops etc.
These are reversible cells; they behave as galvanic cell while discharging and as electrolytic cell while charging.
To improve the performance of battery for commercial purpose
 The anodes and cathodes with very small separation to conserve space are used.
 Current discharge should be high at low temperature.
 It should have less variation in voltage during discharge
 It should be reliable.
 It should have tolerance to shock, temperature etc.
 It should have number of charging and discharging cycles before failure of battery (Cycle life)
P a g e | 13
Lead –acid battery:
If a number of cells are connected in series, the arrangement is called a battery. The lead storage battery is
one of the most common batteries that are used in the automobiles. A 12 V lead storage battery is generally used,
which consists of six cells each providing 2 V. Each cell consists of a lead anode and a grid of lead packed with
lead oxide as the cathode. These electrodes are arranged alternately, separated by a thin wooden piece and
suspended in dil. H2SO4 (38%), which acts as an electrolyte (Fig. 1.13).Hence it is called Lead-acid battery.
Anode: Pb
Cathode:PbO2
Electrolyte: H2SO4 (20.22%)
EMF=2V
To increase the current output of each cell, the cathode and the anode plates are
Joined together, keeping them in alternate positions. The cells are connected parallel
to each other (anode to anode and cathode to cathode). The cell is represented as
Pb | PbSO4 (s), H2SO4 (aq.) | PbSO4 (s), Pb
In the process of discharging, i.e. when battery produces current, the reactions at
the electrodes are as follows:
At anode:
Pb  Pb+2 + 2eAt cathode
Pb+2 (s) + SO42- (aq.)  PbSO4 (s)
PbO2 (s) + SO42- (aq.) + 4H+ (aq.) + 2e–  PbSO4 (s) + 2H2O
Therefore, overall reaction is
Pb (s) + PbO2 (s) + 4H2SO4 (aq.)  2PbSO4 (s) + 2H2O
During discharging the battery, H2SO4 is consumed, and as a result, the density of
H2SO4 falls; when it falls below 1.20 g/cm3, the battery needs recharging. In
Discharging, the cell acts as a voltaic cell where oxidation of lead occurs.
During recharging, the cell is operated like an electrolytic cell, i.e. electrical
energy is supplied to it from an external source. The electrode reactions are the
reverse of those that occur during discharge.
PbSO4 (s) + 2e– Pb (s) + SO4– – (aq.)
PbSO4 (s) + 2H2O  PbO2 (s) + 2H2SO4 + 2e–
2PbSO4 (s) + 2H2O  Pb (s) + PbO2 (s) + 2H2SO4 (aq.)
During this process, lead is deposited at the cathode; PbO2 is formed at the anode
and H2SO4 is regenerated in the cell.
Advantages: Lead acid batteries are used for supplying current to railways, mines, laboratories, hospitals,
automobiles, power stations, telephone exchange, gas engine ignition, Ups (stand-by supplies). Other advantages
are its recharge ability, portability and its relatively constant potential & low cost.
Disadvantages: Use of Conc.H2SO4 is dangerous; Use of lead battery is fragile.
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Nickel–cadmium cell (Ni-cad cell)
It is rechargeable secondary cell. It consists of cadmium as the negative electrode
(anode) and NiO2 acting as a positive electrode (cathode). Potassium hydroxide (KOH) is used
as an electrolyte. The cell reaction during charging and discharging are as follows.
Anode: Cd
Cathode: NiO (OH)
Electrolyte: KOH
EMF=1.4V
At Anode
Cd(S) + 2OH(Aq)  Cd(OH)2 (s) + 2eAt Cathode
NiO(OH) (s) + 2H2O + 2e-  2 Ni(OH)2+ OH-(aq)
Overall reaction
Cd(s) + 2 Ni(OH) + 2H2O  Cd(OH)2 (s) + 2 Ni(OH)2(s)
Advantages and uses
 The Nickel-Cadmium cell has small size and high rate charge/discharge capacity, which makes it very
useful.
 It has also very low internal resistance and wide temperature range (up to 70°C).
 It produces a potential about 1.4 volt and has longer life than lead storage cell.
 These cells are used in electronic calculators, electronic flash units, transistors etc.
 Ni- Cd cells are widely used in medical instrumentation and in emergency lighting, toys etc.
III)
Fuel Cell :
Definition: A Fuel cell is an electrochemical cell which converts chemical energy contained in
readily available fuel oxidant system into electrical energy.
Principle: The basic principle of the fuel cell is same as that of electrochemical cell.
The only difference is that the fuel & oxidant are stored outside the cell. Fuel and Oxidant are supplied
continuously and separately to the electrodes at which they undergo redox reactions.
Fuel cells are capable of supplying current as long as reactants are replenished.
Fuel + Oxidant  Oxidation Products + electricity
Eg : 1)H2 -O2 fuel cell
2) Propane -O2 fuel cell
3) CH3OH-O2 fuel cell
A. Hydrogen – Oxygen fuel cell: One of the most successful fuel cell is H2 –O2 fuel cell.
The cell consists of two inert porous electrodes made of graphite impregnated with finely divided ‘Pt’ (or) Ni
(or) Pd – Ag alloy and a solution of 2.5% KOH as electrolyte.
H2 & O2 gases are bubbled through anode & cathode compartments respectively. The following reactions take
place.
Cell Reaction:
At anode:
2H2 (g) + 4OH-  4H2O + 4eAt Cathode:
O2 (g) + 2H2O + 4e-  4OHNet Reaction:
2H2 (g) + O2 (g)  2H2O, Ecell = 1.23V
A large no of these cells connected in series form a fuel-cell battery. In the production of electricity by this
method, the byproducts are heat, CO2, water, which will not cause pollution of the environment.
P a g e | 15
Applications:
1. These are used as auxiliary energy source in space vehicles, submarines and other military vehicles.
2. The product water produced is a valuable source of fresh water for astronauts.
3. Fuel cell is preferred in spacecraft because of its lightness.
4. Advantages:
5. 1) Fuel cells have high efficiency. It is nearly 70% while other sources have efficiency 15-20% (gasoline
engine) and 30-35 %( diesel engine).
6. 2) The efficiency of the fuel cell does not depend on the size of the power plant.
7. 3) Maintience cost is very low.
8. 4) Fuel cells are more efficient in producing the mechanical power to drive the vehicles and require less
energy consumption.
9. Disadvantages:
10. 1) Initial cost of fuel cell is high.
11. 2) Life time of fuel cell is not known accurately.
12. 3)There is a problem of durability and storage of large amount of hydrogen
Distinction between Primary, Secondary & Fuel cells
Primary
Secondary
1) It only acts as galvanic or 1) It acts as galvanic or
voltaic cell. i.e., produces
voltaic cell while
electricity
discharging and acts as
electrolytic cell
2) Cell reaction is not
2) Cell reaction is
reversible.
reversible.
3) Can’t be recharged
3) Can be recharged
4) Can be used as long as
the active materials are
present
eg: Leclanche cell or Dry
cell, Lithium cell
4) Can be used again and
again by recharging.
eg: Lead storage battery,
Ni-Cd battery, Lithium ion
cell
Fuel cells
1) It is a simple galvanic or
voltaic cell. i.e., produces
electricity
2) Cell reaction is reversible
3) Energy can be withdrawn
continuously
4) Reactants should be
replenished continuously. it
does not store energy
eg: H2&O2 Fuel cell
CH3OH &O2 Fuel cell