This lesson explains how redox reactions work, how electrons move through electrochemical cells, how electrolysis operates, and how electrode potentials determine cell feasibility.

1. Core Concepts (Short Notes)

20.1 Oxidation & Reduction

• Oxidation: Loss of electrons (OIL)

• Reduction: Gain of electrons (RIG)

Oxidizing agents accept electrons. Reducing agents donate electrons.

20.2 Redox Reaction

A reaction in which both oxidation and reduction occur simultaneously. 

Example: Zn → Zn²⁺ + 2e⁻ (oxidation) Cu²⁺ + 2e⁻ → Cu (reduction)

20.3 Electrochemical Cells

Two types:

1. Galvanic/Voltaic cell – generates electricity (spontaneous).

2. Electrolytic cell – requires electricity (non-spontaneous).

2. Galvanic (Voltaic) Cells

20.4 Structure

• Two half-cells connected by a salt bridge.

• External circuit for electron flow.

20.5 Electron & Ion Flow

• Electrons flow anode → cathode.

• Anode = oxidation.

• Cathode = reduction.

• Salt bridge maintains charge balance.

20.6 Cell Potential (E°cell)

Calculated using standard electrode potentials: E°cell = E°cathode − E°anode

If E°cell > 0 → reaction is spontaneous.

Example: Zn/Cu cell → 1.10 V

3. Electrolytic Cells

20.7 Electrolysis

Using electric current to force a non-spontaneous reaction.

20.8 Electrodes & Reactions

In electrolysis:

• Cathode: reduction (cations gain electrons).

• Anode: oxidation (anions lose electrons).

20.9 Examples

Electrolysis of molten NaCl: Cathode: Na⁺ + e⁻ → Na Anode: Cl⁻ → ½Cl₂ + e⁻

Electrolysis of brine (NaCl solution): Cathode: 2H₂O + 2e⁻ → H₂ + 2OH⁻ Anode: 2Cl⁻ → Cl₂ + 2e⁻

4. Faraday’s Laws of Electrolysis

20.10 First Law

Mass deposited ∝ charge passed.

20.11 Second Law

For same charge, mass ∝ equivalent weight.

20.12 Key Formula

Mass = (Q × M) / (nF) Where:

• Q = It (charge)

• M = molar mass

• n = electrons transferred

• F = 96500 C/mol

5. Tips & Tricks for Exams

• Electron flow: anode → cathode (always!).

• Oxidation is loss, reduction is gain.

• In electrolysis of aqueous solutions, water may be reduced/oxidized.

• More reactive metals (e.g., Na, K) are not deposited from aqueous solution.

• Salt bridge must contain inert electrolyte (e.g., KNO₃).

• E°cell > 0 means a feasible cell.

6. Important Points to Remember

• Redox reactions involve electron transfer.

• Galvanic cells convert chemical energy → electrical.

• Electrolytic cells convert electrical → chemical.

• Standard electrode potentials predict spontaneity.

Faraday’s laws relate charge to mass deposited.