CBSE Class 12 Chemistry Revision Notes Chapter 2 Electrochemistry

Electrochemistry studies the relation between chemical reactions and electrical energy.

In CBSE Class 12 Chemistry, this chapter explains cells, electrode potential, conductance, electrolysis, batteries and corrosion.

Electrochemistry deals with two linked processes. Chemical energy can produce electrical energy in a spontaneous redox reaction. Electrical energy can also bring about a non-spontaneous chemical reaction.

Use these CBSE Class 12 Chemistry Revision Notes Chapter 2 to revise the 2026–27 chapter in a quick and exam-ready way. Start with electrochemical cells and Daniell cell. Then revise Nernst equation, conductance, Kohlrausch law, electrolysis, batteries, fuel cells and corrosion.

Key Takeaways

• Electrochemistry: It connects redox reactions with electrical energy.
• Galvanic Cell: It converts chemical energy into electrical energy.
• Electrolytic Cell: It uses electrical energy for a non-spontaneous reaction.
• Nernst Equation: It gives cell potential under non-standard conditions.

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Access 30 Minutes Class 12 Chemistry Chapter 2 Electrochemistry Notes

Electrochemistry can be revised faster if you divide it into cells, formulas and applications. First revise anode, cathode, oxidation and reduction.

Then move to Nernst equation, conductance and electrolysis. Keep batteries, fuel cells and corrosion for final revision because they are more application-based.

Electrochemistry Class 12 Chemistry Chapter 2 Overview

Electrochemistry studies electricity produced from chemical reactions and chemical reactions caused by electricity.

Many industrial processes use electrochemical methods. Metals, sodium hydroxide, chlorine and fluorine can be produced by electrochemical processes.

Batteries and fuel cells convert chemical energy into electrical energy. Corrosion is also explained as an electrochemical process.

Electrochemical Cells and Their Types

An electrochemical cell has two electrodes dipped in electrolytic solution. It uses redox reactions to connect chemical change with electricity.

There are two main types of electrochemical cells.

Galvanic Cell

A galvanic cell converts the chemical energy of a spontaneous redox reaction into electrical energy.

In a galvanic cell:

Electrons flow from anode to cathode through the external circuit. Current flows in the opposite direction.

Electrolytic Cell

An electrolytic cell uses electrical energy to carry out a non-spontaneous redox reaction.

In an electrolytic cell, an external voltage is applied. This forces the reaction to occur in the required direction.

Example: In copper sulphate solution, copper is deposited at the cathode and dissolves at the anode.

Daniell Cell in Electrochemistry Notes

Daniell cell is a galvanic cell made of zinc and copper electrodes dipped in their salt solutions.

Cell reaction:

Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Half-cell reactions:

Zinc acts as anode. Copper acts as cathode.

The cell potential of Daniell cell is 1.1 V when Zn²⁺ and Cu²⁺ ion concentrations are 1 mol dm⁻³.

Cell representation:

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

Galvanic Cell and Electrolytic Cell Difference

In both cells, oxidation occurs at anode and reduction occurs at cathode.

Electrode Potential and Standard Hydrogen Electrode

Electrode potential is the potential difference developed between an electrode and its electrolyte.

When all species in a half-cell have unit concentration, the electrode potential is called standard electrode potential.

The standard hydrogen electrode is used as a reference electrode. Its potential is taken as zero.

Standard hydrogen electrode is represented as:

Pt(s) | H₂(g) | H⁺(aq)

Conditions:

The standard electrode potential of another half-cell is measured with respect to this electrode.

Cell Potential and EMF

Cell potential is the difference between the electrode potentials of cathode and anode.

Ecell = Ecathode - Eanode

For standard conditions:

E°cell = E°cathode - E°anode

If E°cell is positive, the cell reaction is spontaneous.

For Daniell cell:

E°cell = 0.34 V - (-0.76 V)

E°cell = 1.10 V

Nernst Equation and Cell Potential

The Nernst equation gives the electrode potential or cell potential when concentrations are not standard.

For a general cell reaction:

aA + bB → cC + dD

Nernst equation:

Ecell = E°cell - RT/nF loge Q

At 298 K:

Ecell = E°cell - 0.0591/n log Q

Here:

Nernst Equation for Daniell Cell

For Daniell cell:

Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Ecell = E°cell - 0.0591/2 log [Zn²⁺] / [Cu²⁺]

The cell potential increases when Cu²⁺ concentration increases. It decreases when Zn²⁺ concentration increases.

Relation Between Cell Potential and Equilibrium Constant

At equilibrium, Ecell becomes zero.

The relation is:

E°cell = 2.303RT/nF log Kc

At 298 K:

E°cell = 0.0591/n log Kc

This formula helps calculate equilibrium constant from standard cell potential.

Relation Between Cell Potential and Gibbs Energy

The maximum electrical work done by a galvanic cell is linked with Gibbs energy.

ΔrG = -nFEcell

For standard conditions:

ΔrG° = -nFE°cell

If E°cell is positive, ΔrG° is negative. This means the reaction is spontaneous.

Conductance of Electrolytic Solutions

Conductance explains how easily electricity passes through a material or solution.

Metals conduct electricity through electrons. Electrolytic solutions conduct electricity through ions.

Resistance, Conductance, Resistivity and Conductivity

Conductivity depends on the nature of electrolyte, ion size, solvent, concentration and temperature.

Metallic and Electrolytic Conductance

In electrolytic conductance, ions move through solution and reactions occur at electrodes.

Molar Conductivity

Molar conductivity is the conductance of the volume of solution containing one mole of electrolyte.

Λm = κ / c

When κ is in S cm⁻¹ and c is in mol L⁻¹:

Λm = κ × 1000 / c

Unit: S cm² mol⁻¹

Molar conductivity is important because it shows the conducting power of ions produced by one mole of electrolyte.

Conductivity and Molar Conductivity with Dilution

Conductivity and molar conductivity behave differently on dilution.

This is one of the most important points in Electrochemistry Class 12 notes.

Strong Electrolytes

Strong electrolytes dissociate almost completely in solution.

For strong electrolytes, molar conductivity increases slowly with dilution.

Formula:

Λm = Λ°m - A√c

Here:

A graph of Λm against √c gives a straight line for strong electrolytes.

Weak Electrolytes

Weak electrolytes dissociate only partially in solution.

Their molar conductivity increases sharply on dilution. This happens because dilution increases the degree of dissociation.

For weak electrolytes, Λ°m cannot be found by simple extrapolation. Kohlrausch law is used for this.

Kohlrausch Law and Its Applications

Kohlrausch law is also called the law of independent migration of ions.

It states that limiting molar conductivity of an electrolyte is the sum of the individual contributions of its ions.

For electrolyte AxBy:

Λ°m = xλ°+ + yλ°-

Here:

Applications of Kohlrausch Law

Degree of dissociation:

α = Λm / Λ°m

For weak acid:

Ka = cα² / (1 - α)

Electrolytic Cells and Electrolysis

Electrolysis is the process in which electrical energy brings about a chemical change.

In an electrolytic cell, a voltage source is connected to electrodes dipped in an electrolyte.

Example: Electrolysis of copper sulphate solution using copper electrodes.

At cathode:

Cu²⁺(aq) + 2e⁻ → Cu(s)

At anode:

Cu(s) → Cu²⁺(aq) + 2e⁻

Copper dissolves at the anode and deposits at the cathode.

Faraday’s Laws of Electrolysis

Faraday gave two laws to explain quantitative aspects of electrolysis.

First Law of Electrolysis

The amount of chemical reaction at an electrode is proportional to the quantity of electricity passed through the electrolyte.

Q = It

Here:

Second Law of Electrolysis

The amounts of different substances liberated by the same quantity of electricity are proportional to their chemical equivalent weights.

Equivalent weight = Atomic mass / Number of electrons required

One faraday is the charge carried by one mole of electrons.

1 F = 96487 C mol⁻¹

Batteries and Fuel Cells

Batteries are galvanic cells used as sources of electrical energy.

They may be primary batteries or secondary batteries.

Primary Batteries

Primary batteries cannot be recharged after use. Their reactions are not easily reversible.

Examples:

Secondary Batteries

Secondary batteries can be recharged. Their cell reactions can be reversed by passing current.

Examples:

The lead storage battery is a common secondary battery.

Fuel Cells

A fuel cell converts the energy of fuel combustion directly into electrical energy.

Hydrogen-oxygen fuel cell uses hydrogen and oxygen. It produces water as the product.

Overall reaction:

2H₂(g) + O₂(g) → 2H₂O(l)

Fuel cells are considered cleaner because they can produce electricity with less pollution.

Corrosion as an Electrochemical Process

Corrosion is the slow coating of metal surfaces with oxides or other salts.

Rusting of iron is the most common example. It occurs in the presence of air and water.

At anode:

Fe(s) → Fe²⁺(aq) + 2e⁻

At cathode:

O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l)

Overall reaction:

2Fe(s) + O₂(g) + 4H⁺(aq) → 2Fe²⁺(aq) + 2H₂O(l)

Ferrous ions are further oxidised to ferric ions. Rust forms as hydrated ferric oxide.

Rust:

Fe₂O₃.xH₂O

Prevention of Corrosion

Preventing corrosion is important because it protects bridges, buildings, ships and machines.

Quick Formula Table for Electrochemistry

Important Terms in Electrochemistry

Useful Links for Class 12 Chemistry