Electrochemistry

Electrochemistry is that branch of

chemistry which deals with the study
of production of electricity from
energy released during spontaneous
chemical reactions and the use of
electrical energy to bring about non-
spontaneous chemical
transformations.
 Importance of Electrochemistry
1. Production of metals like Na, Mg. Ca and Al.
2. Electroplating.
3. Purification of metals.
4. Batteries and cells used in various
instruments.
 Conductors
Substances that allow electric current to
pass through them are known as
conductors.
Metallic Conductors or Electronic
Conductors
Substances which allow the electric
current to pass through them by the
movement of electrons are called
metallic conductors, e.g. metals.
 Electrolytic Conductors or
Electrolytes
Substances which allow the passage of
electricity through their fused state or
aqueous solution and undergo chemical
decomposition are called electrolytic
conductors, e.g., aqueous solution of
acids. bases and salts.
 Electrolytes are of two types:
1. Strong electrolytes The electrolytes that
completely dissociate or ionise into ions are
called strong electrolytes. e.g., HCl, NaOH, S
2. Weak electrolytes The electrolytes that
dissociate partially (ex < 1) are called weak
electrolytes, e.g., CCOOH, , NOH, S, etc
Electrochemical Cell and Electrolytic
General Representation of an
Electrochemical Cell
Other features of the
electrochemical cell are
1. There is no evolution of heat.
2. The solution remains neutral on both
sides.
3. In the reaction flow of electrons stops
after sometime.
Daniell Cell
 Function of salt bridge
1. It completes the circuit and allows the flow of current.
2. It maintains the electrical neutrality on both sides.
Salt-bridge generally contains solution of strong
electrolyte such as KN, KCl etc. KCI is preferred because
the transport numbers of and are almost same
Transport number or Transference
number
The current flowing through an electrolytic solution
is carried by the ions. The fraction of the current
carried by an ion is called its transport number or
transference number. Thus Transport number of
cation. nc = (current carried by cation/total
current). Transport number of anion na = (current
carried by anion/total current)Evidently nc + na = 1
 Electrode Potential

When an electrode is in contact with the

solution of its ions in a half-cell, it has a
tendency to lose or gain electrons which is
known as electrode potential. It is expressed in
volts. It is an intensive property, i.e.,
independent of the amount of species in the
reaction.
 Oxidation potential

The tendency to lose electrons in the

above case is known as oxidation
potential. Oxidation potential of a
half-cell is inversely proportional to
the concentration of ions in the
solution.
 Reduction potential

The tendency to gain electrons in the

above case is known as reduction
potential. The reduction potential
alone be called as the electrode
potential unless it is specifically
mentioned.
•E°red = – E°oxidation
It is not possible to determine the
absolute value of electrode potential.
For this a reference electrode [NHE or
SHE] is required. The electrode
potential is only the difference of
potentials between two electrodes
that we can measure by combining
them to give a complete cell.
 Standard electrode potential

The potential difference developed

between metal electrode and solution of
ions of unit molarity (1M) at 1 atm
pressure and 25°C (298 K) is called
standard electrode potential.
•It is denoted by E°.
 Reference Electrode

The electrode of known potential is called

reference electrode. It may be primary
reference electrode like hydrogen
electrode or secondary reference
electrode like calomel electrode.
(g) → 2(ag) + 2 Oxidation
(ag) + 2 → Reduction
 Drawbacks of SHE
1. It is difficult to maintain 1 atm pressure of gas.
2. It is difficult to maintain ion concentration 1 M.
3. The platinum electrode is easily poisoned by
traces of impurities.
ELECTROMOTIVE FORCE (EMF)
 Electromotive Force (emf) of a Cell

It is the difference between the electrode

potentials of two half-cells and cause flow
of current from electrode at higher
potential to electrode at lower potential.
It is also the measure of free energy
change. Standard emf of a cell
EQUATION FOR ELECTROCHEMICAL
CELL
Here, ΔG° is the standard Gibbs free energy
change
 Concentration Cells
(i) Electrode concentration cells Two hydrogen
electrodes or different pressures are dipped in
the same solution of electrolyte
(ii) Electrolyte concentration cells : Electrodes are
the same but electrolyte solutions have different
concentrations, e.g..
 Electrochemical Series

It is the arrangement of electrodes in

the increasing order of their standard
reduction potentials at 298 K
 Applications of Electrochemical Series
(ECS)
1. The lower the value of E°, the greater the
tendency to form cation. M→M
+ ne-
Metals placed below hydrogen in ECS
replace hydrogen from dil acids but metals
placed above hydrogen cannot replace
hydrogen from dil acids.
3. Oxides of metals placed below hydrogen are
not reduced by but oxides of iron and metals
placed above iron are reduced by ·
SnO, PbO, CuO are reduced by
CaO, O are not reduced by ·
4. Reducing character increases down the series.
5. Reactivity increases down the series.
6. Determination of emf; emf is the difference of
reduction potentials of two half-cells.
= –
If the value of emf is positive. then reaction take
place spontaneously, otherwise not.
7. Greater the reduction potential of a substance,
oxidising power. (e.g.. > > > )
8. A negative value of standard reduction potential
shows that it is the site of oxidation.
9. Oxides of metals having E°red ≥ 0.79 will be
decomposed by heating to form and metal.
HgO (s) → Hg(l) + (1/2)(g)
(E°/Hg = 0.79V)
 Nernst Equation
The relationship between the concentration of ions and
electrode potential is given by Nernst equation.
For a electrochemical cell,
Concentration of pure solids and liquids is taken as
unity.
Nernst equation and Kc
At equilibrium
Relationship between free energy
change and equilibrium constant
ΔG° = – 2.303RT log Kc
 Conductance (G)

It is the ease of flow of electric current through the

conductor. It is reciprocal of resistance (R)
• G = (1/R), units , mhos or
 Molar Conductivity ()
The conductivity of all the ions produced when 1
mole of an electrolyte is dissolved in V mL of
solution is known as molar conductivity. It is related
to specific conductance as
= (k x 1000/M)
where. M = molarity.
It units are or S .
 Equivalent conductivity ()

The conducting power of all the ions produced

when 1 g-equivalent of an electrolyte is
dissolved in V mL of solution, is called
equivalent conductivity. It is related to specific
conductance as
Λm = (k x 1000/N)
where. N = normality.
 Debye-Huckel Onsagar equation

It gives a relation between molar

conductivity, at a particular concentration
and molar conductivity at infinite dilution.
• = m –B √C
where, B is a constant. It depends upon the
nature of solvent and temperature.
 Factors Affecting Conductivity

(i) Nature of electrolyte The strong electrolytes like

KN KCl. NaOH. etc. are completely ionised in
aqueous solution and have high values of
conductivity (molar as well as equivalent).
The weak electrolytes are ionised to a lesser extent
in aqueous solution and have lower values of
conductivity (molar as well as equivalent) .
ii) Concentration of the solution The concentrated
solutions of strong electrolytes have significant interionic
attractions. which reduce the speed of ions and lower
the value of and .
The dilution decreases such attractions and increase the
value of Λm and .
The limiting value, m or m. (the molar conductivity
at zero concentration (or at infinite dilution) can be
obtained extrapolating the graph.
In case of weak electrolytes, the degree of
ionisation increases dilution which increases the
value of Λm and Λeq. The liminting value mcannot
be obtained by extrapolating the graph.
limiting value, m, for weak electrolytes is obtained
by Kohlrausch law.
(iii) Temperature: The increase of
temperature decreases inter-ionic
attractions and increases kinetic energy
of ions and their speed. Thus, Λm and
Λeq increase with temperature.
 Kohlrausch’s Law
At infinite dilution, the molar conductivity of an
electrolyte is the sum of the ionic conductivities of the
cations and anions, e.g., for AxBy.
 Applications

(i) Determination of equivalent/molar

conductivities of weak electrolytes at infinite
dilution,e.g.,
(ii) Determination of degree of dissociation (α)
of an electrolyte at a given dilution.
The dissociation constant (K) of the weak
electrolyte at concentration C of the solution
can be calculated by using the formula
kc = (C/1 – α)
where, α is the degree of dissociation of the
electrolyte.
(iii) Salts like BaS, PbS, AgCl, AgBr and AgI which do
not dissolve to a large extent in water are called
sparingly soluble salts.
The solubility of a sparingly soluble salt can be
calculated as
 Electrolysis
It is the process of decomposition of an electrolyte when
electric current is passed through either its aqueous
solution or molten state,
1. In electrolytic cell both oxidation and reduction takes
place in the same cell.
2. Anode is positively charged and cathode is negatively
charged, In electrolytic cell.
3. During electrolysis of molten electrolyte, cations are
liberated at cathode. while anions at the anode.
4. When two or more ions compete at the
electrodes. the ion with higher reduction
potential gets liberated at the cathode
while the ion with lower reduction
potential at the anode.
For metals to be deposited on the
cathode during electrolysis, the voltage
required is almost the same as the
standard electrode potential. However
for liberation of gases, some extra voltage
is required than the theoretical value of
the standard electrode potential. The
extra voltage thus required is called over
Discharge potential is defined as the
minimum potential that must be
applied across the electrodes to bring
about the electrolysis and subsequent
discharge of the ion on the electrode.
 Faraday’s Laws of Electrolysis
1. First law
The amount of the substance deposited or liberated at
cathode directly proportional to the quantity of electricity
passed through electrolyte.
W∝Ixt=IxtxZ=QxZ
I current in amp, t = time in sec,
Q = quantity of charge (coulomb)
Z is a constant known as electrochemical equivalent.
When I = 1 amp, t = 1 sec then Q = 1
coulomb, then w = Z.
Thus, electrochemical equivalent I the
amount of the substance deposited or
liberated by passing 1A current for 1 sec
(i.e. 1 coulomb, I x t = Q)
 2. Second law
When the same quantity of electricity is passed through
different electrolytes. the amounts of the substance
deposited or liberated at the electrodes arc directly
proportional to their equivalent weights, Thus,
 Batteries
These are source of electrical energy which
may have one or more cells connected in
series.
 Primary Batteries
In the primary batteries. the reaction
occurs only once and after use over a
period of time battery becomes dead
and cannot be reused again.
(i) Dry cell or Leclanehe cell
• Anode-Zn-Hg amalgam
• Cathode-Paste of (HgO + C)
Electrolyte-Moist paste of KOH-ZnO
 Secondary Batteries

These cells can be recharged and

can be used again and again
• Anode-Spongy lead
• Cathode-Grid of lead packed with Pb
• Electrolyte-38% by mass
•Electrodes-Made of porous graphite
impregnated with catalyst (Pt, Ag or a metal
oxide).
•Electrolyte-Aqueous solution of KOH or NaOH
•Oxygen and hydrogen are continuously fed
into the cell.
 Corrosion
Slow formation of undesirable
compounds such as oxides, sulphides
or carbonates at the surface of metals
by reaction with moisture and other
atmospheric gases is known as
corrosion.
 Factors Affecting Corrosion
1. Reactivity of metals
2. Presence of moisture and atmospheric
gases like C, S, etc.
3. Presence of impurities
4. Strains in the metal
5. Presence of electrolyte
 Rusting of Iron-Electrochemical
Theory
•An electrochemical cell, also known as
corrosion cell, is developed at the surface
of iron.
•Anode- Pure iron
•Cathode-Impure surface
Rusting of iron can be prevented by the following
methods :
1. Barrier protection through coating of paints
or electroplating.
2. Through galvanisation or coating of surface
with tin metal.
3. By the use of antirust solutions (bis phenol).
4. By cathodic protection in which a metal is
protected from corrosion by connecting it to