UNIT- II

Electrochemistry

Prepared by S.I.Davis Presley

 Objective
The students will
• Understand the concept of Redox reaction.
• Conductance and different types of conductance
• Know the origin of Electrode Potential.
• Acquire the knowledge of Electrochemical cell..
• Understand the need for reference electrodes..
• Acquire the skill to calculate electrode potential of half cell.
• To know the principle of Conductometry,pH-metry and
potentiometry.

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 Electrochemistry
• Electrochemistry is the study of chemical processes that
involve electricity caused by movement of electrons and ions.
• Electrochemical reactions are Redox reactions.

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 Redox Reaction
 The term Redox comes from two concepts of reduction and
oxidation.
 Both reduction and oxidation reactions occur together.

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 Redox reaction
• Reduction describes the gain of electrons by acceptor (oxidizing
agent).
Ex., Cu2+ (aq) + 2 e- Cu (s)
• Oxidation describes the loss of electron by an electron donor
(reducing agent).
Ex., Zn(s) Zn2+ (q) + 2 e -
• Each of the reaction is known as half- cell reaction and an electrode
with electrolyte is called half-cell.
• Both half- cell reactions must always go side by side to sustain the
electrochemical reaction.
• The net Redox reaction: Zn (s) + Cu2+(aq) Zn2+ (aq) + Cu (s)

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 Redox reaction
• Reduction describes the gain of electrons by acceptor (oxidizing
agent).
Ex., Cu2+ (aq) + 2 e- Cu (s)
• Oxidation describes the loss of electron by an electron donor
(reducing agent).
Ex., Zn(s) Zn2+ (q) + 2 e -
• Each of the reaction is known as half- cell reaction and an electrode
with electrolyte is called half-cell.
• Both half- cell reactions must always go side by side to sustain the
electrochemical reaction.
• The net Redox reaction: Zn (s) + Cu2+(aq) Zn2+ (aq) + Cu (s)

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 Electrolytes
Any substance that produce ions when
dissolved in a solvent (usually water) is an
electrolyte.
It is the electrically conductive solution
that must be present for corrosion to occur.
Types of electrolytes
Strong electrolyte
Weak electrolyte
Non-electrolyte
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 Strong Electrolytes

Strong electrolytes are substances that only exist as ions in

solution.
They completely dissociate to their ions when dissolved in
solution.
Ionic compounds are typically strong electrolytes.
Strong acids, strong bases and salts are strong electrolytes
They conduct electricity when molten or in aqueous solution.
Example: Hydrochloric acid, Sodium chloride

HCl  H 2O  H 3O   Cl 

 
NaCl  H 2O  Na  Cl
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 Weak Electrolytes

A weak electrolyte only partially dissociates in

solution and produces relatively few ions (exist in
water as a mixture of individual ions and incontact
molecules).
Polar covalent compounds are typically weak
electrolytes.
Weak acids and weak bases are weak electrolytes.
They conduct electricity weakly.
Example: Acetic acid, ammonia
CH 3COOH  H 2O  CH 3COO  H 

9 NH 3  H 2O  NH 4  OH 
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 Non Electrolytes

A non-electrolyte does not dissociate at all (present entirely as

intact molecules) in solution and therefore does not produce
any ions.
Non-electrolytes are typically polar covalent substances that
do dissolve in water as molecules instead of ions.
They do not conduct electricity at all.
Example: Sugar

C12 H 22O11  H 2O  C12 H 22O11

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 Electrical Terms
SI Term SI Symbol SI Unit
Electrical Current I Ampere (A)
Quantity of Electricity Q Coulomb (C)

Electric Potential V Volt(V)

Electric Resistance R Ohm()
Resistivity  m
Conductance G Siemens (S);ohm-1
Conductivity  Sm-1; -1m-1; -1cm-1
Molar conductivity  Sm2mol-1
Molar Conductivity of  Sm2mol-1
Ion
Electric Mobility of u m2V-1s-1
Ion
Transport Number of t
Ion
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 Other Terms

Symbol Term
C Molar Concentration, mol dm-3 (with
:mol m-3)
 Degree of dissociation
l Length
A Area
Kcell Cell Constant
° Molar conductivity at infinite dilution
or Limiting Molar Conductivity

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 Conductance

• What is conductance?
Conductance is an expression of the ease with which electric current flows
through materials like metals and nonmetals.
• There are two types of conductors:
Electronic conductors – flow of electrons (e.g. Metals allows electricity to
pass through them without undergoing any chemical change).
. Electrolytic conductors – flow of ions (e.g. Molten salts, electrolyte
solutions through which electricity is passed through undergoing chemical
changes).
Conductance is represented by G. Conductance in electricity is considered the
opposite of resistance(R ).
It’s a reciprocal relationship between conductance and resistance is expressed
through the following equation

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 Specific Conductance

• Specific conductance is the ability of a substance to conduct electricity. It is

the reciprocal of specific resistance.
• Specific conductance is defined as the conducting capacity of a solution of
the dissolved electrolyte and the whole solution is being placed between two
electrodes are 1sq.cm and length 1cm.
• Specific conductance is denoted by (Kappa).
• It is the reciprocal of resistivity()
• Ohms law states that the resistance of any substance is directly proportional
to its length and inversely proportional to its area. Ohms law is expressed as
•
• ( = Specific resistance)

l = Length,A= Area of Coss Section & R= Resistance

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 • Cell Constant
Note
• When l = 1 Cm and A= 1 Cm
• Then
Hence specific resistance of a conductor can be defined as the resistance of
the conductor having length of 1 cm and area of 1 cm.
Specific conductance (= Kappa)
1/

X 1/R
= Cell constant X conductance(G)

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 Unit of Specific conductance

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 Molar Conductance

• Molar conductivity is the conductance property of a solution containing one

mole of the electrolyte. It is therefore not a constant.

• When one mole of an electrolyte is dissolve in a solution, the ions produced

in the solution are capable of conducting electricity. Molar conductance is
defined as the conducting power of the dissolved ions produced in the
solution.

• It is denoted by (Lambda)

It is connected with the specific conductivity ,  by the relation.

m = /C or ( M is molarity)

Where C is the concentration. The molar conductance is usually expressed in

Scm2 mol-1
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 Equivalent conductance

• It is defined as the conductance of all the ions produced by dissolving one

gram equivalent of an electrolyte in a particular solution.
• We can say that the conductance of an electrolyte solution depends on the
concentration of the ions present in the solution.

• ( N is the Normality)

Factors affecting Electrolytic conductance.

i) Temperature
The conductance of an electrolyte solution increases with increase in
temperature.(1°C increase in temperature causes 2% increase in
conductance)
a) Increases the extent of ionization
b) Decreases the viscosity of the solvent makes ions to move freely to the
electrode.

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 • Concentration of ions- As the number of ions increase, the conductance
increase.
• Nature of Electrolyte
• Ionic Size and Mobility

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 Electrochemistry in our daily life
Corrosion : Decaying of metal and alloys by oxidation half-
cell reaction.

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 Electrochemistry in our daily life
Electroplating: Metal gets deposited at the cathode surface by
reduction half- cell reaction.

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 Electrochemistry in our daily life
Decomposition of ionic substance into simple products by redox
reaction

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 Electrochemical cell
Electrochemical cells are devices that change chemical
energy into electrical energy or change electrical energy
into chemical energy.

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 Parts of electrochemical cell

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 Parts of Electrochemical cell
1. Two half-cells: Each half cell containing an electrode in contact
with an electrolyte. Oxidation occurs at one half cell (anode) while
reduction takes place at the other half cell (cathode).

2. Salt bridge/ Porous Membrane : Both the half-cells are joined by

salt bridge, or some other path (porous membrane or ceramic) that
allows ions to move between the half-cells but prevents mixing of
the electrolytes.

3. External Connection: The two metallic electrodes are joined by

wire through which electrons flow from one half cell to the other
one.

• In some cases both half-cells use the same electrolyte, so that the
electrochemical cell consists of two electrodes in contact with a
single electrolyte.
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 Salt Bridge and its Functions
• Salt bridge contains a solution of a salt inwhich the transport
numbers of anion and cation are the same. ( ex.KCl, KNO 3
NH4NO3).
• It serves as a bridge to complete the electric circuit by joining both
the half-cells.
• It helps to maintain electro neutrality of electrolyte.
• It prevents transference or diffusion of the solutions from one half-
cell to the other.
• It minimizes the liquid junction potential.

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 Liquid Junction potential

 Liquid-liquid junction potential or diffusion potential occurs when two

solutions of different concentrations/composition are in contact with each
other.
 The additional potential develops in the cell due to the difference in
mobilities of anions and cations of the electrolytes across the boundary
between the two solutions.

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 Liquid junction potential in commercial reference electrode

• For precise measurements of potential, use of salt bridge is

preferred.
• The electrode potential measured by commercial reference electrode
uses porous membrane is not accurate due to the development of
additional potential (liquid junction potential) at the porous
membrane.

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 Liquid Junction Potential and Salt Bridge

• The comparable ionic mobility of anion and cation of the salt bridge
minimise the additional potential of opposite sign developed at the
junctions of the salt bridge in both the half-cells.

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 Classification of electrochemical cell
Cell

Electrochemical cell
Electrolytic cell

Galvanic or Voltaic cell

Concentration cell

Chemical cell

Reversible Irreversible Electrolyte or Metal Oxygen Electrode or active

cell cell ion Concentration Concentration Passive Concentration
cell cell cell

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 Classification of Electrochemical Cell

Based on the spontaneity of electrochemical reaction

1.Galvanic or voltaic cell: Produces electrical energy by a
spontaneous chemical reaction involving flow of electrons.
Ex Electrochemical corrosion
2. Electrolytic cell: Stimulates chemical reaction by supplying
electrical current. ex. Electroplating

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 Electrolytic cell Galvanic or Voltaic cell
Converts electrical energy into chemical Converts Chemical energy into electrical
energy energy
The anode carries positive charge vice The anode carries negative charge
versa. Vice versa.
Electrons are supplied to the cell from Electrons are drawn from the cell itself.
an external power supply.
Electrodes usually kept in the same Consists of two separate half cells with
electrolyte. different electrolyte .
The extent of chemical reaction occurring The e.m.f of the cell is theoretically
at the electrode is governed by Faraday’s calculated by Nernst equation.
law of electrolysis.
The amount of electricity passed during The e.m.f produced in the cell is
electrolysis is measured by Coulometer. measured by potentiometer.
Not a spontaneous reaction. Spontaneous reaction.
eg. Electroplating eg. Galvanic Corrosion

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 Classification of Electrochemical Cell Based on
reversibility
1. Reversible Cell: A cell which obey the three conditions of
thermodynamic reversiblity.
• In reversible cell, reactions can be reversed by applying suitable
opposing external emf.
• Ex. Rechargeable batteries.
2. Irreversible Cell: Cell which does not obey the conditions of
thermodynamic reversibility.
• Reactions cannot be reversed by applying emf greater than that of the
cell.
• Ex. Primary batteries.

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 Types of galvanic cells

Based on the way electricity is produced

1) Chemical cell: Electricity is produced due to potential gradient developed

at both half cells caused by redox reactions..
Ex. Daniel cell.

2) Concentration Cell: Electricity is produced due to the movement of ion by

concentration gradient.
Ex. Differential Aeration corrosion

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 Representation of an Anode half cell
• The anodic cell is represented by writing the metal first and then the
metal ions present in the electrolyte. These two are separated by a
vertical line or a semicolon.
• For example:

• The molar concentration or activity of the solution is written in

brackets after the formula of the ion.

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 Representation of Cathode half cell

• The cathode half cell is represented by writing the metal ion first of
the electrolyte solution and then the metal. Both are separated by
vertical line or semicolon.
• For example:

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 Representation of electrochemical cell
• Anode is written on the left hand side and cathode on the right hand
side.
• Salt bridge is represented by two vertical lines.

• So Daniel cell is represented as:

• If the concentration of both the electrolytes is 1 M, then the cell

notation is:

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 Electromotive force (e.m.f )of the cell
• Difference of potential which causes the flow of current from the electrode
of higher potential to one of lower potential is called the “electromotive
force(emf)” of the cell.
• It is the algebraic sum of the single electrode potentials provided the proper
sign are given according to actual reaction taking place on the electrodes.

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 Points to remember
• Electrochemical cells are classified into different types based on
spontaneity and reversibility of electrochemical chemical reaction as
well as the way of production of electricity.
• Emf of the cell is defined as the difference of potential which causes
the flow of current from the electrode of higher potential to one of
lower potential.
• Emf of the cell can be theoretically calculated by Nernst equation
and measured by potentiometer-using Poggendorff’s compensation
principle(wheatstone's bridge).

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 Origin of Electrode Potential
• When a metal(M) is placed in the solution of its own ion it may
lose e-s (act as an anode) or gain e-s (act as a cathode) depends
on its electron affinity.
1. Oxidation: A metal atom on the electrode M may lose n electrons to
the electrode, and enter the solution as Mn+, (i.e., the metal atom is
oxidised).

2. Reduction: Movement of positive ions from the solution to the metal

electrode to get reduced .

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 Development of Negative charge at the metal surface
by Oxidation

 The electrons left by the metal on  The negatively charged electrode attracts
oxidation get accumulated on the positively charged free ions (cations) from
metal electrode. the solution which remain close to the
 The electrode acquires a slight metal.
negative charge with respect to the  A region of charge unbalance is created all
solution . around the metal.

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 Accumulation of positive charge at the surface by
Reduction

 Movement of positive ions from the  Accumulated of positive charge at the

solution to the metal electrode in electrode attracts negatively charged
order to get reduced (deposited) on ions resulting in the formation of region
the electrode . of charge unbalance around the metal.

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 Accumulation of Charges at the surface of metal electrodes

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 Formation of Helmholtz Double Layer
• The regions of electrical unbalance (positive or negative ions)
formed all around the metal when the metal is placed in the
solution of its own ion is known as Helmholtz Double Layer.

• Formation of this layer prevents further passing or deposition

of metal ions across the layer, leads to establishment of
equilibrium.

• At equilibrium, the potential difference set up between metal

and electrolyte is constant that is known as the electrode
potential.

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 Helmholtz double layer and electrode
potential
The electrode
potential is a
measure of
tendency of an
electrode to lose or
gain electron.

Inner Helmholtz Layer (compact Layer)

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 Oxidation/Reduction potential
• Oxidation Potential: The tendency of an electrode to lose electrons
or to get oxidized is called its oxidation potential.
• Reduction Potential: The tendency of an electrode to gain electrons
or to get reduced is called its reduction potential.
• The oxidation potential is the reverse of reduction potential.

• For example, reduction potential of Zn is -0.76 volts, its oxidation

potential is +0.76 volts.

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 Standard electrode potential

• Standard electrode potential is a measure of tendency of metal

to lose or gain electron when the pure metal is in contact with
its own ions at one molar concentration at a temperature of
25oC or 298 K.

• When a Zn rod is dipped in 1M ZnSO4 solution, standard electrode is

formed and the potential developed is called standard zinc electrode
potential (Eo Zn+2/Zn).

• The standard zinc electrode is represented as Zn+2(IM)/Zn.

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 Standard Electrode potential of gas Electrode
• The standard electrode potential (Eo) of gas electrode is defined as
the potential developed at the interface of the gas and solution
containing its own ions of unit concentration at a pressure of 760 mm
of Hg or 1atm.

• When the H2 gas at a pressure of 1atm is bubbled through 1M HCl

solution, std H2 electrode is formed and the potential developed is
called std hydrogen electrode potential (E oH2) whose magnitude is
considered to be 0.

• The standard H2 electrode is represented as Pt, H2 / H+(760 mm of

Hg)/ (IM)

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 Types of Electrodes
Electrode type Example Description Electrode reaction (in reduction
direction)

Metal metal-ion Cu(s)│Cu2+(aq) Metal immsersed in electrolyte Cu2+(aq) +2e ⇌ Cu(s)

electrode containing its own ions.

Ion – ion (redox) Pt(s)│Fe3+,Fe2+(aq) Noble metal in contact with solution Fe3+(aq)+e⇌Fe2+(aq)
electrode of a redox couple

Metal insoluble salt Hg(s) │Hg2Cl2(s) │KCl(aq) Metal in contact with its sparingly Hg2Cl2(s)+2e ⇌ 2Hg+2Cl-
electrode soluble salt and a solution
containing a soluble anion of the
salt

Gas electrode Pt(s)│H2(g) │H+(aq) Noble metal in contact with a H+(aq)+e ⇌1/2H2(g)
saturated solution for a gas and
contains the reduced or oxidized
form of the gas.

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 Theoretical Calculation of electrode potential

Nernst Equation
• In an electrode reaction, the decrease in free energy produces electrical
energy: ΔG° = -nFE°
• For a reversible reaction, Mn+(aq) + ne - ⇌ M(s)
• ΔG = ΔG° + RT ln [product]/[reactant]---- Von't Hoff Reaction
isotherm
-nFE = -nFE° + RT ln [product]/[reactant]

• E=E° + ln[reactant]/[product] = E° + ln[Mn+]/ [M]

• As the concentration of the solid M is taken as unity and substituting the

values of R, F and T=298 K (Room Temperature)
• Electrode Potential, E = E° + ln[Mn+]

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 Factors influencing Electrode potential

The electrode potential depends upon

• the concentrations of the substances or pressure in
the case of a gas electrode.
• the temperature.

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 Measurement of Electrode potential

• The absolute value of a single electrode potential cannot be

measured experimentally because a half-cell reaction cannot
take place independently.
• But its value in relation to the potentials of other half cells
(reference electrodes) can be measured.
• Its unit is volt (V).

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 Reference Electrodes
• Reference electrode is a non-polarizable electrode with constant
potential which acts as one of the half cells of the electrochemical
cell to measure the electrode potential of an electrode as a cell.
Types of Reference electrodes
1. Primary Reference electrode
• The electrode whose potential is arbitrarily fixed to a certain value
(zero volt).
• Standard Hydrogen Electrode(SHE) or Normal Hydrogen Electrode
(NHE).
2. Secondary reference electrodes
• The potentials of these electrodes are precisely known with respect to
the SHE. Ex. Calomel electrode
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 Requisites for reference electrodes

• Electrodes should be reversible and obey the Nernst equation with

respect to any one of the ions in the electrolyte.
• Easy to prepare and maintain.
• It should act as either anode or cathode.
• Potential should be stable. i.e., its potential should return to the
equilibrium potential after small currents are passed through the
electrode .
• Temperature coefficient of electrode potential should be negligible.

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 Why SHE is chosen as primary reference
electrode?
• Small potential is developed on the hydrogen electrode, hence it can
be taken as zero.

• In determining the single electrode potential, using S.H.E. as a

reference, the potential of the unknown potential will be equal to the
e.m.f. of the cell.

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 Standard Hydrogen electrode
Construction
Hydrogen gas at 1 atm pressure is allowed to bubble over a platinum
electrode having a specially treated (platinised) surface which is
dipped in a 1M H+ solution at 25°
•Platinum electrode catalyzes the reaction,
H+ + e– ⇌ ½ H2 (g)
Electrode :Pt-H2
Electrolyte : HCl
Electrode Representation : Pt,H2(1atm)/H+(1M)

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 Why is SHE seldom used?
• Construction is difficult.
• The impurities present in H2 and HCl poison the Pt, and affect the
equilibrium at the electrode.
• It is difficult to get pure, dry hydrogen gas and prepare ideal
platinised platinum plate.
• The platinum surface has to be specially treated by preliminary
electrolysis.
• The supply of hydrogen gas makes it cumbersome and hazardous.
• It is difficult to maintain the pressure of hydrogen gas.
• It requires large volume of test solution.

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 Standard Calomel Electrode

Type- Metal-metal Insoluble salt electrode

Electrode- Pt-Hg
Electrolyte –Hg2Cl2(s)/ Satd.KCl
Electrode representation Pt-Hg/ Hg2Cl2(s)-KCl(Sat. Solutin)

Hg2Cl2(g

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 Construction

• The calomel electrode consists of a broad glass tube

having sidearm as shown. The sidearm is used for
dipping it any solution used for coupling the calomel
electrode. At the bottom of the glass tube there is pure
mercury and the platinum wire is sealed into the bottom
for electrical connections. The wire runs through the
separator glass tube to the top of the tube for electrical
contact. Above pure mercury there is a paste of
Mercurous chloride(Hg2Cl2) in mercury. The rest of the
glass vessel is filled with 0.1M or 1M KCl solution.

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 Reactions

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 • The electrode potential of calomel electrode varies with the concentration of
KCl as shown below.

Type of Calomel Decinormal (0.1)N Normal (1N KCl) Saturated KCl

Electrode KCl

Potential (V) 0.3335 0.2810 0.2422

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 Ion Selective Electrodes

• An electrode that responds to a particular ions activity is called ion selective

electrode.(ISE).ISE allows measured ions to pass but excludes the passage
of other ions.

• Principle: Whenever two solutions of different concentrations are separated

by a membrane, a potential difference is set up.

Nernst Equation for glass electrode

n=1 and T= 25° C

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 Measurement of Electrode potential
• To measure the relative potential of an electrode, it is coupled with
SHE to form an electrochemical cell and the cell potential is
measured with the help of potentiometer.

• According to modern convention, if on coupling with SHE reduction

occurs at the given electrode, the electrode potential is given positive
sign.

• The potential value is given negative sign if oxidation occurs at the

electrode.

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 Measurement of electrode potential
• The cell potential of
Mg2+ (aq) | Mg(s)|| H+ || H2 (g)| Pt
is - 2.37V
• The electrode potential of Mg2+| Mg is - 2.37V as electrode potential of SHE is zero.
• The negative sign indicates oxidation occurs at the electrode.

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 Electrochemical series
• The arrangement of metals in the increasing order
(downwards) of their standard (reduction) electrode
potential.

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 Electrochemical series -Applications
Relative ease of oxidation or reduction
• The metals on the top in the electrochemical series are more active,
easily ionized into solution and hence they are good electron donor
while the metals at the bottom easily get reduced by accepting
electron and act as good reducing agents.
• Ex., a very high negative reduction potential of lithium electrode
indicates that it is very difficult to reduce Li + ions to Li atoms. But
Li+ loses electrons easily and behaves as a better reducing agent.
• Metals at the bottom of the table have high reduction potential and
they can be easily reduced. Therefore, they act as strong oxidizing
agents. F2 is a better oxidizing agent than Cl2.

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 Electrochemical series -Applications
Corrosion Tendency of Metals
• The chemical reactivity of metals decreases from top to bottom in the
series.
• The metal higher in the series is more active than the metal lower in
the series.
• Alkali metals and alkaline earth metals having high negative values
of standard reduction potentials and are chemically active. These
metals react even with cold water and evolve hydrogen.
• Metals like Cu, Ag and Au which lie below hydrogen are less
reactive do not react with water even at high temperature.

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 Electrochemical series -Applications
Liberation of Hydrogen in Acids
• All metals having negative electrode potentials show greater
tendency of losing electrons as compared to hydrogen. So, when
such a metal is placed in an acid solution, the metal gets oxidized,
and H+ (hydrogen) ions get reduced to form hydrogen gas.

• Zn (E (Zn2+ /Zn) = - 0.76 V), Iron (E (Fe2+ /Fe) = - 0.44 V) etc., can displace
hydrogen from acids.

• But metals such as Ag (E (Ag+ /Ag) = + 0.80V) and Au (E (Au3+ Au) =

+1.42 V) cannot displace hydrogen from acids because of their
positive reduction potential value.

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 Electrochemical series -Applications
Displacement tendency
• A metal occupying higher position in the series can displace the
metals lying below it from the solutions of their salts.
• For example, Zn lies above Cu in the electrochemical series,
therefore, if Zn metal strip is immersed in CuSO 4 solution, Cu is
displaced from the solution.

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 Electrochemical series -Applications
Calculation of Std emf of the cell
Electromotive force (emf) of the cell
 EMF of the cell is the difference of potential that drives the flow of
current from electrode with high potential to the electrode with
lower potential.
 Std emf of the cell is calculated from the standard potential of both
electrodes by using the formula

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 EMF calculation
Std emf of the cell, Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
is calculated as
•Cu2+ + 2 e - ⇌ Cu ------ E°c = +0.337
Zn2+ + 2 e - ⇌ Zn------- E0a = - 0.763
•Eocell = E0c –E0a
•Eocell = 0.337 - (- 0.763) = 1.100 V

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 Electrochemical series -Applications

Predicting the spontaneity/direction of redox reaction

• Spontaneity of the redox reaction can be predicted from the
e.m.f value of complete cell reaction.
• If the value of E cell is +ve , the reaction is feasible.
• If it is -ve the reaction is not feasible.

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 E0 = E0C- E0a = + 1.50 –(-0.74) =+ 2.24

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 Conductometry

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 Strong Acid Vs Strong Base

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 Mixture of Acids Vs Strong Base

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 Precipitation – Conductometric Titration

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 Potentiometry

• Potentiometry is the field of electroanalytical chemistry in which the potential

is measured in an electrochemical cell. Potential is measured under the
conditions of no current flow. The measured potential is proportional to the
concentration of some come component of the analyte.Generally
concentration of the analyte.
• A typical electrochemicalcell for potentiometric studies is constructed by
coupling two half-cells. One is a reference electrode whose reaction and
potential is constant and the other is known as the indicator electrode.
• Any redox reaction can be studied using a potentiometer. The most common
system is the ferrous-ferric system. Ferrous ion can be oxidized to Ferric ion
by an oxidizing agent eg. KMnO4 and K2Cr2O7.
• Generally saturated calomel electrode is coupled with inert platinum
electrode which is the indicator electrode of the cell during titration.

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 Points to remember
• Electrochemical reactions are redox reaction in which both reduction
and oxidation reactions occur together.
• Standard electrode potential is a measure of tendency of metal to lose or
gain electron when the pure metal is in contact with its own ions at one
molar concentration at a temperature of 25oC or 298 K.
• The absolute value of a single electrode potential cannot be measured
experimentally. But its relative value can be measured by coupling
with reference electrode.
• Electrode potential can be theoretically calculated by Nernst
equation.
• Reference electrodes are used to measure the electrode potential.

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 Anna University questions
1. Define the term single electrode potential? (May 2011)
2. What is salt bridge? Explain its functions (June 2011)
3. What is electrode potential? How is it developed? (May 2009)
4. Define single electrode potential. Mention the factors affecting it
(June 2010, June 20130
5. What is electrochemical series? What is its significance? (May
1998, Jan 2013)
6. What are redox reaction, illustrate with an example. (May 2004)
7. Define the term single electrode potential. Derive Nernst equation.
Explain the various terms involved. (May 2009, Jan 2003)

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