CBSE Class 12 Chemistry Chapter 3 (Electrochemistry) By Homi Institute

1. CBSE Class 12 Chemistry
Electrochemistry

2. Electrical Conductivity
It is the measure of a material's ability to
allow the transport of an electric charge.

3. Electrochemistry is the study of chemical processes that cause electrons to move. This movement
of electrons is called electricity, which can be generated by movements of electrons from one
element to another in a reaction known as an oxidation-reduction ("redox") reaction.

7. What is a Salt Bridge?
It is a device which is use in an
electrochemical cell for connecting
its oxidation and reduction half cells,
wherein a weak electrolyte is used. It
usually consists of a strong
electrolyte.
Example, AgNO3, KCl, etc.
The main function of a salt bridge is to help maintain the electrical neutrality within the internal circuit. It
also helps in preventing the cell from taking its reaction to equilibrium.
If salt bridges are absent or if they are not used then the reaction will likely continue and the solution in one-
half electrodes will gather a negative charge. Similarly, in the other half, electrodes would accumulate a
positive charge. This will further result in the stoppage of the reaction and no electricity will be produced.
In case of Electrolysis too, just like any electrochemical cell, the salt bridge will have the same function.
*Note – if two electrodes are in one solution in a single container, no salt bridge is necessary.

8. Glass tube salt bridge
Types of Salt Bridges
There are mainly two types
of salt bridges used in
electrochemical cells.
• Glass Tube Bridge
• Filter Paper/Tissue Paper
Bridge
A salt bridge basically helps in preventing the accumulation of positive and negative charges
around the respective electrodes and further allowing a smooth reaction to take place. It also helps
in the continual flow of electrons. However, the purpose of a salt bridge is not to move electrons
from the electrolyte; rather it’s to maintain charge balance because the electrons are moving from
one-half cell to the other.
• Salt bridge prevents the diffusion or mechanical flow of solution from one-half cell to another.
• It prevents or minimizes the liquid-liquid junction potential. (Potential arises between two
solutions when they are in contact with each other).
• Salt bridge acts as an electrical contact between two half cells.

9. For Example
• Oxidation half cell (anode)

10. Also known as Single Electrode Potential
According to IUPAC convention, standard reduction
potentials are now called standard electrode potentials.

11. It is the most commonly used reference electrode since its potential is exactly equal to Zero at all
temperatures. Since the potential of SHE is “zero”, it forms the perfect basis to calculate cell
potentials using different electrodes or different concentrations.
Disadvantages of Standard Hydrogen Electrode (S.H.E.):
1. It is not convenient to assemble the apparatus.
2. It is difficult to maintain the pressure of hydrogen gas at 1atm and concentration of HCl at 1M
all the time.
3. It is difficult to get pure, dry hydrogen gas and prepare ideal platinised platinum plate.
4. Its very costly and is difficult to transport due to its bigger size.

12. Determination of Standard Electrode Potential of Zn Electrode
To determine the Standard Electrode Potential
of Zn 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.
Potential difference = 1.10V

13. Determination of Standard Electrode Potential of Cu Electrode
To determine the Standard Electrode Potential of Cu
electrode, a copper rod is dipped in 1 M solution of
copper sulphate solution. It is combined with hydrogen
electrode through a salt bridge. Both the electrodes joined
through a voltmeter. The deflection of the voltmeter
indicates that current is flowing from copper electrode
towards hydrogen electrode, i.e., the electrons are moving
from hydrogen to copper electrode. The hydrogen
electrode acts as an anode and the copper electrode as
a cathode.
Potential difference = 1.10V

15. Nernst equation & Gibbs free energy
The Nernst Equation enables the
determination of cell potential under
non-standard conditions. It relates the
measured cell potential to the reaction
quotient and allows the accurate
determination of equilibrium constants
(including solubility constants).
The Nernst Equation is derived from the Gibbs free energy under standard conditions.
(1)
ΔG is also related to E under general conditions (standard or not) via
(2)
with
• n is the number of electrons transferred in the reaction (from balanced reaction),
• F is the Faraday constant (96,500 C/mol), and
• E is potential difference.

16. Under standard conditions, Equation 2 is then
(3)
Hence, when E⁰ is positive, the reaction is spontaneous and when E⁰ is negative, the reaction is non-spontaneous.
From thermodynamics, the Gibbs energy change under non-standard conditions can be related to the Gibbs energy
change under standard Equations via
(4)
into Equation 4, we have:
(5)
Divide both sides of the Equation above by −nF, we have
(6)
Equation 6 can be rewritten in the form of
(7)
and Equation 7 can be rewritten as:
(8)
The Equation above indicates that the electrical potential of a cell depends upon the reaction quotient Q of the reaction. As the
redox reaction proceeds, reactants are consumed, and thus concentration of reactants decreases. Conversely, the products
concentration increases due to the increased in products formation. As this happens, cell potential gradually decreases until the
reaction is at equilibrium, at which ΔG=0 . At equilibrium, the reaction quotient Q=Keq. Also, at equilibrium, ΔG=0 and
ΔG=−nFE, so E=0.

17. Therefore, substituting Q=Keq and E=0 into the Nernst Equation, we have:
(9)
At room temperature, Equation 9 simplifies into (notice natural log was converted to log base 10):
(10)
This can be rearranged into:
(11)
The Equation above indicates that the equilibrium constant Keq is proportional to the standard potential of the
reaction. Specifically, when:
• K>1,E⁰ >0 , reaction favours products formation.
• K<1,E⁰ <0, reaction favours reactants formation.
This result fits Le Châtlier's Principle, which states that when
a system at equilibrium experiences a change, the system will
minimize that change by shifting the equilibrium in the
opposite direction.

18. Example
The E⁰cell = +1.10V for the Zn-Cu redox reaction:
(12)
What is the equilibrium constant for this reversible reaction?
Answer -

19. Conductance of Electrolytic Solutions

22. Conductivity Cell
A conductivity cell is composed of two platinum electrodes
which are coated with platinum black. The electrodes have
area of cross section equal to A and are separated by distance
‘l’. Hence, the solution confined between the two electrodes is
a column of length l and area of cross section A. The
resistance for this column of solution is given by following
equation.

28. Batteries

29. Mercury Cell
Dry Cell Button Cell
• Dry Cell
• Mercury Cell
• Lead Acid Cell
• Ni-Cd Cell

30. Anode
Cathode

31. Ni-Cd Cell

33. Oxidation
Reduction
Bisphenol Chemical is toxic for Environment
Tin (Sn) Coating
Mg as a sacrificial coating