CBSE Important Questions Class 12 Chemistry Chapter 2 – Electrochemistry

Class 12 Chemistry Chapter 2 Electrochemistry is part of the CBSE Class 12 Chemistry syllabus. It deals with the relationship between chemical reactions and electrical energy. Students learn about galvanic cells, standard electrode potential, the Nernst equation, conductivity of electrolytic solutions, Kohlrausch's law, electrolysis and Faraday's laws, and the working of batteries, fuel cells and corrosion.

These CBSE important questions for Class 12 Chemistry Chapter 2 cover Faraday's laws applied to electrolysis, the SI unit of resistance, the standard hydrogen electrode, a case study on the dry cell and lead versus cadmium batteries, fuel cells, identifying anode and cathode in a given cell, and why solid NaCl does not conduct electricity. Every question has a clear step-by-step solution in simple English with figures, so students can practise on their own, parents can check the working, and teachers can use them for class tests.

Q1. A solution is electrolysed for 'x' minutes with a current of 1.6 A to deposit 0.3482 g of iron. Find the value of x. [1 F = 96500 C, atomic mass of Fe = 56 g]

Answer: Iron is deposited from Fe²⁺, so n = 2 electrons are needed per atom.
Equivalent mass of Fe = 56/2 = 28 g
By Faraday's first law, w = (M × I × t) / (n × F)
0.3482 = (56 × 1.6 × t) / (2 × 96500)
0.3482 × 193000 = 89.6 t
67202.6 = 89.6 t
t = 750 seconds
x = 750/60 = 12.5 minutes, that is about 13 minutes

Q2. Deduce the units of resistance based on the SI convention.

Answer: From Ohm's law, R = V/I.
Potential difference V = work done per unit charge = joule/coulomb
R = (J/C) / A
Now J = kg m² s⁻², and C = A s.
R = (kg m² s⁻²) / (A s × A)
= kg m² s⁻³ A⁻²
This combination of base units is given the name ohm (Ω).

Q3. The standard reduction potential of the hydrogen electrode is
(a) +0.30 V
(b) −0.58 V
(c) +0.58 V
(d) 0 V

Answer: (d) 0 V
The standard hydrogen electrode (SHE) is taken as the reference electrode and its potential is defined as exactly zero at all temperatures.
All other electrode potentials are measured against it. The conditions are 1 bar pressure of H₂ gas, 1 M H⁺ ion concentration and a platinum electrode.

Q4. Read the paragraph and answer the questions that follow.
A battery is a device made of one or more electrochemical cells that convert chemical energy into electrical energy. In a dry cell the electrolyte is immobilised as a paste. The zinc-carbon battery (Leclanché cell) has an outer zinc container as the anode, and the cathode is a central carbon rod surrounded by a carbon and manganese dioxide mixture. A paste of ammonium chloride is the electrolyte.
(i) What is the oxidation state of manganese after the reaction in a dry cell?
(ii) Which ion or atom is formed at the anode?
(iii) Mohit is shown two car batteries, one with cheaper lead plates and one with costlier cadmium plates. The shopkeeper suggests the cadmium one. Do you agree? Give two reasons.

Answer: (i) At the cathode, MnO₂ is reduced to MnO(OH). The oxidation state of manganese changes from +4 to +3.
(ii) The zinc container is the anode. Zinc is oxidised, so Zn²⁺ ions are formed:
Zn → Zn²⁺ + 2e⁻
(iii) Yes, the shopkeeper is right. Two reasons:
1. Lead is toxic to humans and causes serious health problems, especially in children. Cadmium batteries (nickel-cadmium) avoid this particular hazard in use.
2. Longer life and better performance — nickel-cadmium cells have a much longer working life than lead storage batteries and do not need the maintenance that lead-acid cells do.
The value here is choosing a long-term benefit to health and the environment over a short-term saving in money.

Q5. (a) How do fuel cells produce electricity? (b) What are the electrode reactions in a hydrogen-oxygen fuel cell?

Answer: (a) A fuel cell converts the chemical energy released by the combustion of a fuel directly into electrical energy, without actually burning it. The fuel and the oxidant are supplied continuously from outside, so the cell keeps working as long as they are fed in.
(b) Electrode reactions in a H₂-O₂ fuel cell:
At the anode: 2H₂(g) + 4OH⁻(aq) → 4H₂O(l) + 4e⁻
At the cathode: O₂(g) + 2H₂O(l) + 4e⁻ → 4OH⁻(aq)
Overall: 2H₂(g) + O₂(g) → 2H₂O(l)
The only product is water, which is why fuel cells are considered clean. Their efficiency is about 70%, much higher than a thermal power plant.

Q6. Consider the electrochemical cell shown in the figure. (a) Identify the reactions at the anode and cathode. (b) Represent the cell.

Diagram of a galvanic cell with two electrodes connected by a salt bridge

Answer: In a galvanic cell, oxidation always happens at the anode and reduction at the cathode. The metal with the more negative electrode potential acts as the anode.
(a) At the anode (oxidation): Zn(s) → Zn²⁺(aq) + 2e⁻
At the cathode (reduction): Cu²⁺(aq) + 2e⁻ → Cu(s)
(b) Cell representation: The anode is written on the left and the cathode on the right, with a double vertical line for the salt bridge:
Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

Q7. In an electrochemical cell
(a) kinetic energy changes into potential energy
(b) chemical energy changes into electrical energy
(c) potential energy changes into kinetic energy
(d) kinetic energy changes into chemical energy

Answer: (b) Chemical energy changes into electrical energy
This is the definition of a galvanic (voltaic) cell. A spontaneous redox reaction is used to drive electrons through an external circuit, producing electricity.
The reverse happens in an electrolytic cell, where electrical energy is used to drive a non-spontaneous chemical reaction.

Q8. Solid sodium chloride does not conduct electricity due to
(a) the dative nature of the compound
(b) repulsion of ions because of the applied electrical force
(c) strong electrostatic forces of attraction holding the ions in place
(d) the covalent nature of the compound

Answer: (c) Strong electrostatic forces of attraction holding the ions in place
In solid NaCl the Na⁺ and Cl⁻ ions are locked in a rigid crystal lattice by strong electrostatic forces. Since the ions cannot move, there are no mobile charge carriers and no current flows.

Q9. (a) Solid NaCl does not conduct electricity but aqueous NaCl does. Give the reason. (b) Give two examples of strong electrolytes.

Answer: (a) In solid NaCl the ions are held fixed in the lattice by strong electrostatic forces, so they cannot move and cannot carry current.
When NaCl is dissolved in water, the high dielectric constant of water weakens the attraction between Na⁺ and Cl⁻. The ions separate and become free to move. These free ions carry the current, so the aqueous solution conducts electricity.
Molten NaCl also conducts for the same reason.
(b) Two strong electrolytes: sodium chloride (NaCl) and hydrochloric acid (HCl). Other examples are H₂SO₄, NaOH and KCl. A strong electrolyte ionises almost completely in solution.

Q10. Write the Nernst equation for an electrode reaction and state what it is used for.

Answer: For the electrode reaction Mn+ + ne⁻ → M, the Nernst equation is
E = E° − (0.059/n) log (1 / [Mn+])  at 298 K
For a full cell it is written as
Ecell = E°cell − (0.059/n) log Q
What it is used for: It lets us calculate the electrode or cell potential under non-standard conditions, that is when the concentrations are not 1 M. It also shows how the cell voltage falls as the reaction proceeds and the products build up.

Q11. State Kohlrausch's law of independent migration of ions and give one of its applications.

Answer: Kohlrausch's law: The limiting molar conductivity of an electrolyte is the sum of the individual contributions of its cation and anion.
Λ°m = ν₊ λ°+ + ν₋ λ°−
Application: It is used to find the limiting molar conductivity of a weak electrolyte, such as acetic acid, which cannot be measured directly by extrapolation because its conductivity does not reach a limiting value at low concentration.
For example, Λ°(CH₃COOH) = Λ°(CH₃COONa) + Λ°(HCl) − Λ°(NaCl).
It is also used to calculate the degree of dissociation and the dissociation constant of a weak electrolyte.

Q12. What is corrosion? Write the reactions involved in the rusting of iron and give two methods of preventing it.

Answer: Corrosion is the slow eating away of a metal by its reaction with the surrounding air and moisture. The rusting of iron is the most common example and is an electrochemical process.
At the anode (oxidation): 2Fe(s) → 2Fe²⁺ + 4e⁻
At the cathode (reduction): O₂(g) + 4H⁺ + 4e⁻ → 2H₂O(l)
The Fe²⁺ is then further oxidised by atmospheric oxygen to Fe³⁺, which forms hydrated ferric oxide, Fe₂O₃·xH₂O — this is rust.
Two methods of prevention:
1. Barrier protection: covering the surface with paint, grease or a layer of another metal so that air and moisture cannot reach the iron.
2. Galvanisation / cathodic protection: coating the iron with zinc, or connecting it to a more reactive metal such as magnesium. The more reactive metal corrodes instead, protecting the iron.

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