Electrochemistry: Complete Chapter Revision for Class 12 NEET by PW

Electrochemistry connects chemical reactions with electrical energy and is an important part of Class 12 Chemistry. The chapter covers electrolysis, galvanic cells and conductance, with concepts such as electrode reactions, oxidation, reduction, cell potential and the relationship between chemical energy and electrical energy.

PW offers Electrochemistry Class 12 notes where major concepts are arranged topic-wise with important equations, comparisons and numerical applications. Electrolytic cells, galvanic cells, Faraday’s laws, salt bridges, EMF, the Nernst equation, equilibrium constants and concentration cells are also covered for Class 12 and NEET preparation.

Three Major Parts of Electrochemistry

Electrochemistry can be broadly divided into three sections:

The major conceptual focus of the chapter is on electrolytic cells and galvanic cells, along with the quantitative relationships used in electrochemical calculations.

Electrolytic Cells and Galvanic Cells

The fundamental difference is the spontaneity of the chemical reaction.

A phone battery behaves as a galvanic cell during use because its spontaneous chemical reaction produces electricity. During charging, external electrical energy drives the reverse reaction, so the battery behaves as an electrolytic cell.

Anode and Cathode

The anode and cathode are identified by the type of reaction occurring at each electrode, not simply by their charge.

Thus:

Anode = Oxidation

Cathode = Reduction

These definitions apply to both galvanic and electrolytic cells.

The charge of each electrode depends on the type of cell:

For an aqueous sodium chloride electrolytic cell, hydrogen is formed at the cathode and chlorine can be formed at the anode under the usual conditions considered in electrolysis.

Redox reactions occur at the electrodes, while electrons travel through the external conducting circuit.

Electrolysis and Faraday’s Laws

Electrolysis calculations can be approached through the sequence:

Current and time → Charge → Faradays → Moles of electrons → Moles of substance

The charge passed through an electrolytic cell is: Q = It

where:

• Q = charge

• I = current

• t = time in seconds

One mole of electrons carries approximately: 1 Faraday = 96500 C

Therefore: Moles of electrons = It/96500

Selectivity in Aqueous Electrolysis

An aqueous electrolyte contains ions from both the dissolved substance and water. For example, aqueous sodium chloride may contain:

• Na⁺

• Cl⁻

• H⁺

• OH⁻

These ions do not necessarily discharge simultaneously. The preferential discharge depends on their relative electrode potentials and the conditions of electrolysis.

Cation Selectivity

Between sodium ions and hydrogen ions, hydrogen is preferentially discharged under the usual aqueous electrolysis conditions: H⁺ + e⁻ → 1/2 H₂

Anion Selectivity

In aqueous sodium chloride under the usual electrolysis conditions, chloride ions are oxidised at the anode: Cl⁻ → 1/2 Cl₂ + e⁻

For aqueous sodium chloride or lithium chloride:

• Hydrogen forms at the cathode.

• Chlorine forms at the anode under the stated conditions.

• Sodium hydroxide or lithium hydroxide remains in solution.

• The solution becomes more basic.

Inert and Reactive Electrodes

An inert electrode provides a conducting surface without participating chemically in the overall reaction. Platinum and graphite are common examples.

A reactive electrode participates in the electrode reaction.

At a reactive metal anode: M → Mⁿ⁺ + ne⁻

For copper sulfate solution with copper electrodes:

Cathode: Cu²⁺ + 2e⁻ → Cu

Anode: Cu → Cu²⁺ + 2e⁻

Copper is deposited at the cathode while the copper anode dissolves.

A mercury cathode provides a special case in aqueous sodium chloride, where sodium can be discharged and form a sodium amalgam with mercury.

Faraday’s First and Second Laws

Faraday’s First Law

Faraday’s First Law states that the mass of a substance deposited or liberated during electrolysis is directly proportional to the charge passed.

W ∝ Q

Therefore: W = ZQ = ZIt

where Z is the electrochemical equivalent.

Z = Mass deposited/Charge passed

Faraday’s Second Law

Faraday’s Second Law applies when electrolytic cells are connected in series. The same current and therefore the same quantity of charge pass through each cell.

Thus: W₁/W₂ = E₁/E₂

where E₁ and E₂ represent the respective equivalent masses.

Therefore, the masses deposited are proportional to their equivalent masses.

Galvanic Cells and Salt Bridges

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

In a zinc-copper cell:

• Anode oxidation: Zn → Zn²⁺ + 2e⁻

• Cathode reduction: Cu²⁺ + 2e⁻ → Cu

The electrons flow through the external wire from zinc to copper.

Functions of a Salt Bridge

A salt bridge:

• Maintains electrical neutrality in the half-cells.

• Completes the internal electrical circuit.

• Allows ion migration between the half-cells without directly mixing their solutions.

• Generally contains ions with similar ionic mobility.

• Should contain ions that do not react significantly with the half-cell solutions.

For a KCl salt bridge:

• Cl⁻ moves towards the zinc compartment.

• K⁺ moves towards the copper compartment.

Cell Notation

Cell notation provides a compact representation of a galvanic cell.

The basic arrangement is: Anode | Anode solution || Cathode solution | Cathode

Important conventions include:

• Left side represents the anode and oxidation.

• Right side represents the cathode and reduction.

• A single vertical line represents a phase boundary.

• A double vertical line represents the salt bridge.

A metal-metal ion half-cell can be written as: Metal | Metal ion

A gas-ion half-cell requires an inert electrode: Inert electrode | Gas | Ion

A redox half-cell containing only ionic species also requires an inert electrode: Inert electrode | Oxidised ion, Reduced ion

The electrode reactions should always be considered along with the cell notation.

EMF, Gibbs Energy and Standard Potential

The relationship between Gibbs energy and cell EMF is: ΔG = −nFE

For a spontaneous galvanic reaction:

• ΔG < 0

• E > 0

For a non-spontaneous reaction under the specified direction:

• ΔG > 0

• E < 0

Under standard conditions, the cell potential is represented by E°.

The standard hydrogen electrode (SHE) is assigned: E° = 0 V

EMF is an intensive property. Multiplying a balanced chemical equation does not multiply the EMF, while reversing the reaction changes the sign of the EMF.

The relationship between oxidation and reduction potentials is: Oxidation potential = −Reduction potential

A higher reduction potential indicates a greater tendency for reduction and generally corresponds to stronger oxidising character.

A higher oxidation potential indicates a greater tendency for oxidation and generally corresponds to stronger reducing character.

Nernst Equation and Reaction Quotient

The Nernst equation gives the cell potential under non-standard conditions: E = E° − RT ln Q/nF

At 298 K: E = E° − 0.0591 log Q/n

The reaction quotient Q contains the concentrations and gas pressures of the reacting species raised to their stoichiometric powers. Pure solids and pure liquids are not included.

For the zinc-copper cell: Zn + Cu²⁺ → Zn²⁺ + Cu

If: [Zn²⁺] = 0.1 M

and

[Cu²⁺] = 0.01 M

Then: Q = [Zn²⁺]/[Cu²⁺] = 0.1/0.01 = 10

Using: E° = 1.10 V And n = 2

we get: E = 1.10 − (0.0591 × log 10)/2

Therefore: E ≈ 1.07 V

Equilibrium Constant and Cell Potential

At equilibrium:

• E = 0

• ΔG = 0

• Q = K

The relationship between standard cell potential and equilibrium constant at 298 K is:

E° = 0.0591 log K/n

Therefore:

• E° > 0 → K > 1

• E° < 0 → K < 1

• E° = 0 → K = 1

This relation connects the electrical tendency of a cell reaction with its equilibrium position.

Concentration Cells

A concentration cell contains the same electrode material at both electrodes but different concentrations of the corresponding electrolyte.

Examples include:

• Zinc-zinc concentration cell

• Copper-copper concentration cell

A zinc-copper cell is not a concentration cell because its two electrodes are made of different materials.

For a concentration cell:

E°cell = 0

The EMF arises from the difference in concentration between the two half-cells.

The Electrochemistry chapter connects redox reactions with electrical energy through electrolysis, galvanic cells and electrochemical calculations. These PW Electrochemistry Class 12 notes bring together electrode reactions, Faraday’s laws, salt bridges, cell notation, EMF, the Nernst equation, equilibrium constants and concentration cells in a topic-wise format. 

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