ELECTROCHEMISTRY

Why Study Electrochemistry?

It is important because
• Many of devices that we
use every day that are
battery powered.
• A big problem called corrosion
is an electrochemical phenomenon.
• Many chemicals and elements
such as Cl2, NaOH, Al etc.
are produced electrolytically.
 DEFINITION

It is the study of the inter-conversion

between chemical energy and
electrical energy.
 ELECTROCHEMICAL CELLS
GALVANIC CELL

What is an Electrochemical Cell?

An electrochemical cell is a device that can
generate electrical energy from the chemical
reactions occurring in it, or use the electrical
energy supplied to it to facilitate chemical reactions
in it. These devices are capable of converting
chemical energy into electrical energy, or vice
versa.
Such cells capable of generating an electric
current from the chemical reactions occurring in
them care called Galvanic cells or Voltaic cells.
Alternatively, the cells which cause chemical
reactions to occur in them when an electric current
is passed through them are called electrolytic cells.
 ELECTRO CHEMICAL CELLS
A COMPARISON

GALVANIC CELL ELECTROLYTIC CELL

Chemical energy is transformed into electrical Electrical energy is transformed into chemical
energy in these electrochemical cells. energy in these cells.
The redox reactions that take place in these An input of energy is required for the redox
cells are spontaneous in nature. reactions to proceed in these cells, i.e. the
reactions are non-spontaneous.

In these electrochemical cells, the anode is These cells feature a positively charged anode
negatively charged and the cathode is and a negatively charged cathode.
positively charged.

The electrons originate from the species that Electrons originate from an external source
undergoes oxidation. (such as a battery).
 DANIEL CELL
In Daniel’s cell, copper ions are reduced at the cathode while
zinc is oxidized at the anode.
Reactions of Daniel cell at cathode and anode are:

At cathode: Cu 2+ + 2e– → Cu

At anode: Zn → Zn2+ + 2e–

 ELECTRODE POTENTIAL
 Following two changes occur when a metal rod is dipped
in its salt solution, 
(a) Oxidation: Metal ions pass from the electrode into
solution leaving an excess of electrons and thus a
negative charge on the electrode.
The conversion of metal atoms into metal ions by the
attractive force of polar water molecules.
M →  Mn + ne-
(b) Reduction: Metal ions in solution gain electrons from
the electrode leaving a positive charge on the
electrode. Metal ions start depositing on the metal
surface leading to a positive charge on the metal.
Mn+ +  ne- →  M
In the beginning, both these changes occur
with different speeds but soon an
equilibrium is established.

M Mn+ + ne-
In practice, one effect is greater than the
other,
If first effect is greater than the second, the
metal acquires a negative charge with
respect to solution and
If the second is greater than the first, it
acquires positive charge with respect to
solution, thus in both the cases a potential
difference is set up.
 REFERENCE ELECTRODE
• A reference electrode is an electrode which has a
stable and well-known electrode potential.
• There are many ways reference electrodes are used.
The simplest is when the reference electrode is used
as a half-cell to build an electrochemical cell. This
allows the potential of the other half cell to be
determined. An accurate and practical method to
measure an electrode's potential in isolation
(absolute electrode potential) has yet to be
developed.
 STANDARD HYDROGEN
ELECTRODE (SHE)
It is reference electrode consists of a
platinum electrode in contact with H2
gas (1 atm) and aqueous H+ ions (1 M).
It is assigned 0.0 V electrode potential.
It may behave as anodic or cathodic
half cell.
It is represented as
Pt(s)|H2(g)(aH = 1)|H+(aq)(aH+ = 1).
2

When SHE is coupled with an other

half cell then cell potential is the value
of the electrode potential of half cell.
 Determination of Standard Electrode Potential of
Zn/Zn2+ Electrode

•A zinc rod is dipped in 1 M zinc sulphate solution. This half-

cell is combined with a standard hydrogen electrode through
a salt bridge.
•Both the electrodes are con­nected with a voltmeter.
•The deflection of the voltmeter indicates that current is
flowing from hydrogen electrode to metal electrode or the
electrons are moving from zinc rod to hydrogen electrode.
•The zinc electrode acts as an anode and the hydrogen
electrode as cathode Oxidation
Cell Representation and the half
cellreaction
can be represented as     
Reduction half reaction

Zn|Zn2+ (aq)/Anode(-) Zn →  Zn2+ + 2e- 2H+ + 2e- →  H2↑

|| 2H(aq)|
H2 (g)/Cathode (+)
•The EMF of the cell is 0.76 volt
E0CELL =E0CATHODE - E0ANODE

0.76 =  EoAnode + 0
EoAnode = +0.76 V
 NERNST EQUATION
Reduction Potential under Non-standard Conditions is
determined using Nernst Equation when Concentrations
not-equal to 1M.

where:
E = actual ½ cell reduction potential

2.303RT [M ] Eo = standard ½ cell reduction potential

E  Eo  log 
nF [M ] n = number of electrons in reaction

0.059
o [1] T = temperature (K)
EE  log 
n [M ] R = ideal gas constant (8.314J/(K-mol)
F = Faraday’s constant (96500 C/mol)
 Nernst Equation for a cell
Reduction Potential under Non-standard Conditions is determined using
Nernst Equation when Concentrations not-equal to 1M. Thus For the cell,
Cu(s)I Cu2+(aq)IIZn2+ (aq)IZn(s)
With cell reaction. Cu2+ (aq) + Zn(s) Zn2+ (aq) +Cu(s)

2
o 2 . 303 RT [ Zn ]
EE  log
nF [Cu 2 ]
2
o 0.059 [ Zn ]
EE  log
2 [Cu 2  ]
 NERNST EQUATION......
c d
0. 05916 [C ] [ D ]
Ecell  Eo  log a b
n [ A] [ B]

At equilibrium Ecell =0:

o 0.05916
E cell  log K
n
The value of E˚cell is related to Gibbs free energy, ∆G˚ by:

∆G˚ = –nFE˚cell

The value of Equilibrium constant, K is related to ∆G˚ by:

∆G˚ = –2.303RT logK
 EQUILIBRIUM CONSTANT AND GIBB’S
FREE ENERGY FROM NERNST EQUATION
At equilibrium Ecell = 0

There fore E0cell = 0.059/n log Kc

log Kc= E0cellX n/0.059

0rG = – nFE0(cell)

Where 0rG is standard Gibb’s free energy of reaction

and E0(cell) is standard emf of the cell
 RESISTENCE & CONDUCTIVITY
• Resistance, R, is proportional to the distance, l, between the electrodes and is
inversely proportional to the cross-sectional area of the sample, A (noted S on the
Figure above). Writing ρ (rho) for the specific resistance (or resistivity),

R =l/Aρ
• In practice the conductivity cell is calibrated by using solutions of known specific
resistance, ρ*, so the quantities l and A need not be known precisely. If the
resistance of the calibration solution is R*, a cell-constant, C, is derived
R*= C X ρ*
• The specific conductance (conductivity), κ (kappa) is the reciprocal of the specific
resistance.
κ = 1/ρ = C/R
• Conductivity is also temperature-dependent. Sometimes the ratio of l and A is
called as the cell constant, denoted as G*, and conductance is denoted as G. Then
the specific conductance κ (kappa), can be more conveniently written as
κ = G* X G
 MOLAR CONDUCTIVITY
• The conductivity is defined as the conductivity of an
electrolyte solution divided by molar concentration
• Molar conductivity = λm = κ/c
• if k is expressed in Sm–1 and the concentration, c in mol
m–3 then the unit of λm will be Sm2mol–1.
• If we use Scm–1 as the units for k and mol cm–3, the units
of concentration, then the units for λm are Scm2mol–1. It
can be calculated by using the equation:
• λm (Scm2mol–1) = κ(S cm–1) × 1000 (cm3/L)/molarity
(mol/L).
• OR, 1 Sm2mol–1 = 104 Scm2mol–1
 CONDUCTIVITY CELL
 While measuring the
resistance (DC) changes
the composition of the
solution. So we use an
alternating current (AC)
source.

 A solution cannot be
connected to the bridge like
a metallic wire, so we use a
specially designed vessel
called conductivity cell.
 VARIATION OF CONDUCTIVITY & MOLAR
CONDUCTIVITY WITH DILUTION
• Conductivity decreases
with dilution
• Molar conductivity
increases with dilution
For strong electrolytes

Λm=Λ0m-AC1/2
 KOHLRAUSCH’S LAW
 It states that limiting molar conductivity of an
electrolyte is the sum of the individual contributions
of the anion and cation of the electrolyte.

 Thus, if λ°Na+ and λ°Cl - are limiting molar

conductivity of the sodium and chloride ions , then
the limiting molar conductivity for sodium chloride
is given by the equation:
Ë°m(NaCl) = λ°Na+ + λ°Cl - .
 APPLICATION OF KOHLRAUSCH’S LAW

• This law may be used to determine the

limiting molar conductivity, ‘λ°m’ degree of
dissociation ‘α’ and dissociation constant
‘Ka’ of a weak electrolyte.
 ELECTROLYTIC CELL AND
ELECTROLYSIS
 Electrolysis: It is the process in which electrical energy is used to
drive a non-spontaneous chemical reaction.

 An electrolytic cell is an apparatus for carrying out electrolysis.

 Processes in an electrolytic cell are the reverse of those in a

galvanic cell.

 Electrolysis process is used in Manufacture of Cl 2 and NaOH,

Electro-refining and Electroplating, Electrolysis of water etc.
 FARADAY’S LAW OF ELECTROLYSIS
 (i) First Law: The amount of chemical reaction which occurs at
any electrode during electrolysis by a current is proportional to
the quantity of electricity passed through the electrolyte (solution
or melt).
 (ii)Second Law: The amounts of different substances liberated
by the same quantity of electricity passing through the
electrolytic solution are proportional to their chemical equivalent
weights (Atomic Mass of Metal ÷ Number of electrons required
to reduce the cation).
 BATTERIES

 Batteries are the most important practical

application of galvanic cells.
 Single-cell batteries consist of one galvanic cell.
 Multi-cell batteries consist of several galvanic
cells linked in series to obtain the desired
voltage.
 Types of batteries

PRIMARY BATTERY SECONDARY BATTERY

In these batteries, the  A secondary battery after

use can be reused by
reaction occurs only once
recharging by passing
and after use over a time
current through it in the
period battery becomes
opposite direction
dead and cannot be
 Example: Lead storage cell
reused again.
 Example: Dry Cell
 DRY CELL
A dry cell is a type of electric battery, commonly used for
portable electrical devices. It was developed in 1886 by the
German scientist Carl Gassner, after development of wet zinc-
carbon batteries by Georges Leclanché in 1866. The modern
version was developed by Japanese Yai Sakizo in 1887.
A dry cell uses a paste electrolyte, with only
enough moisture to allow current to flow
A common dry cell is the zinc-carbon cell, sometimes called
the dry Leclanché cell, with a nominal voltage of 1.5 volts, the
same as the alkaline cell.
It is also called Leclanche cell

Anode: Zinc metal can

Cathode: MnO2 and carbon paste

Electrolyte: NH4Cl and ZnCl2 paste.

Cell Potential: 1.5 V but decreases

to 0.8 V with use.

Anode: Zn(s)  Zn2+ (aq) + 2e-

Cathode: 2NH+4 (aq) + 2MnO2 (s) + 2e-  Mn2O3 (s) + 2NH3 (aq) + H2O (l)

Overall: Zn(s) + 2NH+4(aq) + 2MnO2(s)  Zn2+ (aq) + 2NH3(aq) + H2O(l) + Mn2O3(s)

 MERCURY CELL/ BUTTON CELL
A mercury battery (also called mercuric oxide
battery, mercury cell, button cell, or Ruben-Mallory) is a
non-rechargeable electrochemical battery, a primary cell.

Mercury batteries use a reaction between mercuric oxide

and zinc electrodes in an alkaline electrolyte.

The voltage during discharge remains practically constant at

1.35 volts, and the capacity is much greater than that of a
similarly sized zinc carbon battery. Mercury batteries were
used in the shape of button cells for watches, hearing aids,
cameras and calculators, and in larger forms for other
applications
Anode: Zn(Hg) + 2OH- (aq)  ZnO(s) + H2O (l) + 2e-

Cathode: HgO(s) + H2O (l) + 2e-  Hg(l) + 2OH- (aq)

Overall: Zn(Hg) + HgO(s)  ZnO(s) + Hg (l)

 LEAD STORAGE CELL

Anode : Pb(s) + SO2- (aq) PbSO4(s) + 2e-

Cathode:
PbO2(s) + 4H+ (aq) + SO42- (aq) + 2e- PbSO4(s) + 2H2O(l)
Overall reaction:
Pb(s) + PbO2(s) + 4H+(aq) + 2SO42-(aq) 2PbSO4 (s) + 2H2O (l)
 NICKEL CADMIUM BATTERY
• The nickel–cadmium battery (Ni-Cd battery or NiCad battery) is a type
of rechargeable battery using nickel oxide hydroxide and
metallic cadmium as electrodes.
A fully charged Ni-Cd cell contains:
• a nickel(III) oxide-hydroxide positive electrode plate
• a cadmium negative electrode plate
• an alkaline electrolyte (potassium hydroxide).
• anode reaction
• (i) Cd(s) + 2OH-(aq) ==> Cd(OH)2(s) + 2e-
cathode reaction
• C
• It is rechargeable
• Cell Potential :1.30 V
• Electrolyte: NiO(OH).

Anode: Cd(s) + 2OH–(aq)  Cd(OH)2(s) + 2e–

cathode reaction
(ii) NiO2(s) + 2e-  + 2H2O ==> Ni(OH)2(s) + 2OH-(aq)
overall cell reaction
(iii) Cd(s) + NiO2(s) + 2H2O ==> Cd(OH)2(s) + Ni(OH)2(s)
 FUEL CELL
It’s a type of galvanic cell that requires a continuous supply of
reactants to keep functioning.
Fuel cells are not batteries because they are not self-contained.

It uses externally fed CH4 , CH3OH or H2, which react to form water.

Electrolyte: Hot aqueous KOH solution.

Cell Potential: 1.23 V and have about 40% conversion to electricity;

the remainder is lost as heat. Excess heat can be used to drive
turbine generators.
Anode : 2H2 (g) + 4OH- (aq) 4H2O (l) + 4e-
Cathode: O2 (g) + 2H2O (l) + 4e- 4OH- (aq)
Overall: 2H2 (g) + O2 (g) 2H2O (l)
 CORROSION
 Corrosion is the oxidative deterioration of metal.

 25% of steel produced goes to replace steel

structures and products destroyed by corrosion.

 Rusting of iron requires the presence of both

oxygen and water.

 Rusting results from tiny galvanic cells formed

by water droplets.
Oxidation: Fe(s)  Fe2+(aq) + 2 e–
Reduction: O2(g) + 4 H+(aq) + 4 e–  2 H2O(l)
Overall: 2 Fe(s) + O2(g) + 4 H+(aq)  2 Fe2+(aq) + 2 H2O(l)
 Prevention of corrosion
 Galvanizing: is the coating of iron with zinc. Zinc is more easily
oxidized than iron, which protects and reverses oxidation of the
iron.
 Cathodic Protection: is the protection of a metal from corrosion
by connecting it to a metal (a sacrificial anode e.g. Mg or Zn) that
is more easily oxidized.
 Electroplating.
 By applying paint, grease, rubber to prevent contact of metal
surface from air.