JEE Main Chemistry Redox Reactions & Electrochemistry 2027: Nernst Equation, Cells & Faraday's Laws

A battery is just a controlled redox reaction. Electrochemistry connects redox chemistry to thermodynamics and to everything from phone batteries to metal plating, and the chapter's JEE Main questions revolve around a small set of formulas.

E° tells you which way electrons want to flow. The Nernst equation tells you what happens when concentrations change.

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Chapter at a Glance

Priority reflects past-paper trends, not an official NTA weightage.

What the NTA Syllabus Covers

• Oxidation and reduction, oxidation number, rules for assigning oxidation numbers, balancing redox reactions
• Electrolytic and metallic conduction, conductance, molar conductivity, Kohlrausch's law
• Electrochemical cells, electrode potentials, EMF, Nernst equation, relation between cell potential and Gibbs energy
• Dry cell, lead accumulator, fuel cells

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Master These Topics

1. Cell EMF, ΔG and the Nernst Equation

For a galvanic cell, E°cell = E°cathode − E°anode (both as reduction potentials). Oxidation happens at the anode and reduction at the cathode.

The link to thermodynamics is ΔG° = −nFE°, with F = 96500 C/mol.

Worked example (Daniell cell): For Zn and Cu, E° = 0.34 − (−0.76) = 1.10 V. Then ΔG° = −2 × 96500 × 1.10 ≈ −212 kJ/mol.

Under non-standard conditions the Nernst equation at 298 K is:

E = E° − (0.0591 / n) log Q

Worked example: With [Zn²⁺] = 0.1 M and [Cu²⁺] = 1 M, Q = 0.1, so E = 1.10 − (0.0591/2) × log 0.1 = 1.10 + 0.0296 ≈ 1.13 V. Lowering the product concentration raises the EMF. At equilibrium E = 0, and log K = nE° / 0.0591.

Trap: Do not multiply E° by the stoichiometric coefficient when balancing half-reactions. E° is intensive, while ΔG° depends on n.

2. Conductance and Kohlrausch's Law

Conductivity is κ, and molar conductivity is Λm = κ × 1000 / C when κ is in S cm⁻¹ and C in mol/L. For strong electrolytes, Λm rises slowly on dilution, while for weak electrolytes it rises sharply because ionisation increases.

Kohlrausch's law says that at infinite dilution, Λm° = λ°(cation) + λ°(anion). Worked example: λ°(Na⁺) = 50.1 and λ°(Cl⁻) = 76.3 S cm² mol⁻¹, so Λm°(NaCl) = 126.4. It is also used to find Λm° of weak electrolytes indirectly and their degree of dissociation α = Λm / Λm°.

3. Faraday's Laws of Electrolysis

Mass deposited is m = (M / nF) × I t.

Worked example: A current of 2 A for 965 s passes 1930 C, which is 0.02 mol of electrons. Copper (Cu²⁺ + 2e⁻ → Cu) needs 2 mol of electrons per mole, so 0.01 mol is deposited: 0.01 × 63.5 = 0.635 g.

4. Oxidation Numbers and Batteries

Oxidation numbers: Cr in K₂Cr₂O₇ is +6, and Mn in KMnO₄ is +7. Balance redox reactions using the ion-electron (half-reaction) method.

• Lead accumulator: discharge reaction is Pb + PbO₂ + 2H₂SO₄ → 2PbSO₄ + 2H₂O, and it is recharged by reversing it.
• Fuel cell: H₂ and O₂ react to give water with high efficiency and no harmful emissions.

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Common Traps to Avoid

• Reversing the sign convention: E°cell uses reduction potentials for both electrodes.
• Multiplying E° by coefficients while balancing the cell reaction.
• Forgetting the factor n in the Nernst equation and in ΔG° = −nFE°.
• Using the wrong unit for κ when computing molar conductivity.

60-Second Revision Sheet

• E°cell = E°cathode − E°anode; ΔG° = −nFE°
• Nernst at 298 K: E = E° − (0.0591/n) log Q
• Λm = κ × 1000 / C; Λm° = λ⁺ + λ⁻
• Faraday: m = (M/nF) I t, F = 96500 C/mol

Your Study Plan

1. Day 1: oxidation numbers and balancing redox reactions.
2. Day 2: cell EMF, electrode potentials and ΔG relations.
3. Day 3: Nernst equation with concentration cells.
4. Day 4: conductance, Kohlrausch and Faraday's laws, timed set.

Practice Redox Reactions & Electrochemistry Questions Free → (opens in a new tab)

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Continue Your Chemistry Journey

• Previous chapter: Equilibrium
• Next chapter: Chemical Kinetics
• All JEE Main Chemistry chapters
• Complete JEE Main Syllabus 2027 guide

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