Electrochemistry Class 12 Notes (2026-27) — CBSE
Class 12 Chemistry Chapter 2 notes: galvanic cells, electrode potential, Nernst equation, conductivity, Kohlrausch's law, electrolysis and batteries.
Electrochemistry — Class 12 Chemistry Notes
Chapter Snapshot
Electrochemistry links chemical reactions and electricity. It covers galvanic cells (chemical → electrical), electrode potentials and cell EMF, the Nernst equation for non-standard conditions, the relation to Gibbs energy, conductivity and Kohlrausch's law, Faraday's laws of electrolysis, and practical cells — batteries, fuel cells, and corrosion.
Board relevance: ~9 marks with Chemical Kinetics. Expect a Nernst-equation numerical and a conductivity/Kohlrausch question. Getting the anode/cathode convention right is half the battle.
Key Concepts & Definitions
Galvanic (voltaic) cell — converts chemical energy into electrical energy via a spontaneous redox reaction (e.g. the Daniell cell: Zn Zn²⁺ Cu²⁺ Cu).
Electrolytic cell — uses electrical energy to drive a non-spontaneous reaction.
Anode Cathode
Reaction Oxidation Reduction
Sign (galvanic) Negative (−) Positive (+)
Sign (electrolytic) Positive (+) Negative (−)
Cell notation: anode on the left, cathode on the right —
Zn(s) Zn²⁺(aq) ‖ Cu²⁺(aq) Cu(s)
The salt bridge completes the circuit and maintains electrical neutrality.
Standard electrode potential (E°) — measured against the Standard Hydrogen Electrode (SHE), whose potential is defined as 0.00 V. A more positive E° means a stronger tendency to be reduced (a better oxidising agent).
Formulas — EMF, Nernst and Thermodynamics
Cell EMF:
E°cell = E°cathode − E°anode (both as reduction potentials)
A positive E°cell ⇒ spontaneous reaction.
Nernst equation (at 298 K):
E = E° − (0.059/n) log Q
For the cell reaction aA + bB → cC + dD, Q = [C]^c[D]^d / [A]^a[B]^b. For a single electrode: E = E° − (0.059/n) log(1/[Mⁿ⁺]).
Thermodynamic links:
ΔG° = −nFE° (F = 96500 C/mol)
ΔG° = −2.303 RT log K, so E° = (0.059/n) log K at 298 K
At equilibrium, Ecell = 0 and Q = K.
Conductance
Quantity Definition Unit
Resistance R R = ρ·l/A ohm (Ω)
Conductivity κ κ = 1/ρ = (1/R)(l/A) S m⁻¹
Cell constant G = l/A m⁻¹
Molar conductivity Λm Λm = κ × 1000/M S cm² mol⁻¹
Variation with dilution:
- Both κ decreases on dilution (fewer ions per unit volume) while Λm increases.
- Strong electrolytes: Λm rises slowly and linearly with √C, extrapolating to Λm° (Debye–Hückel–Onsager: Λm = Λm° − A√C).
- Weak electrolytes: Λm rises steeply near infinite dilution (dissociation increases), so Λm° cannot be found by extrapolation.
Kohlrausch's law of independent migration of ions:
Λm° = ν₊λ°₊ + ν₋λ°₋
The limiting molar conductivity is the sum of the individual ionic contributions. Uses: find Λm° for weak electrolytes, and calculate the degree of dissociation α = Λm/Λm° and the dissociation constant Ka = Cα²/(1 − α).
Electrolysis and Practical Cells
Faraday's laws of electrolysis:
1. The mass deposited is proportional to the charge passed: m = Z·I·t (Z = electrochemical equivalent).
2. For the same charge, masses deposited are proportional to their equivalent masses.
m = (M × I × t)/(n × F), F = 96500 C/mol
Practical cells:
- Primary (non-rechargeable): dry cell (Leclanché), mercury cell — used where a steady voltage is needed.
- Secondary (rechargeable): lead storage battery (Pb anode, PbO₂ cathode, H₂SO₄), nickel–cadmium cell.
- Fuel cell (e.g. H₂–O₂): converts the energy of combustion directly into electricity, with high efficiency (~70%) and water as the only product — used in the Apollo space programme.
Corrosion — an electrochemical process; rusting of iron forms hydrated Fe₂O₃. Prevention: painting, greasing, galvanising (zinc coating) and cathodic protection with a more reactive sacrificial metal (Mg, Zn).
Worked Examples
Example 1 — Cell EMF: For Zn Zn²⁺ Cu²⁺ Cu, E°(Zn²⁺/Zn) = −0.76 V and E°(Cu²⁺/Cu) = +0.34 V. Find E°cell.
E°cell = E°cathode − E°anode = 0.34 − (−0.76) = +1.10 V (positive → spontaneous).
Example 2 — Gibbs energy: Find ΔG° for the above cell (n = 2).
ΔG° = −nFE° = −2 × 96500 × 1.10 = −212300 J = −212.3 kJ/mol.
Example 3 — Nernst: Find E for the same cell when [Zn²⁺] = 0.1 M and [Cu²⁺] = 0.01 M.
Q = [Zn²⁺]/[Cu²⁺] = 0.1/0.01 = 10. E = 1.10 − (0.059/2)log 10 = 1.10 − 0.0295 = 1.0705 V.
Example 4 — Molar conductivity: κ = 0.0248 S/cm for a 0.2 M solution. Find Λm.
Λm = κ × 1000/M = 0.0248 × 1000/0.2 = 124 S cm² mol⁻¹.
Example 5 — Faraday: What mass of copper (M = 63.5, n = 2) deposits when 2 A flows for 965 s?
m = (63.5 × 2 × 965)/(2 × 96500) = 122555/193000 ≈ 0.635 g.
Example 6 — Equilibrium constant: For a cell with E° = 0.46 V and n = 2, find log K at 298 K.
E° = (0.059/n) log K → 0.46 = (0.059/2) log K → log K = 0.46 × 2/0.059 ≈ 15.6, so K ≈ 4 × 10¹⁵ — a very large value, confirming the reaction goes essentially to completion. This "convert E° into K" step is a favourite three-marker.
Important Question Patterns
1. Cell EMF & spontaneity (2–3 marks): E°cell = E°cathode − E°anode; predict spontaneity; write cell notation and electrode reactions.
2. Nernst equation (3 marks): compute E at given concentrations; find concentration from a given E; E at equilibrium.
3. Thermodynamics (2–3 marks): ΔG° = −nFE°; relate E° to the equilibrium constant K.
4. Conductivity (3 marks): κ and Λm calculations; how each varies with dilution; strong vs weak electrolyte graphs; Kohlrausch's law and degree of dissociation.
5. Electrolysis/cells (2–3 marks): Faraday's laws numerical; lead storage battery and fuel cell reactions; corrosion and its prevention.
⚡ Quick Revision
- Anode = oxidation (− in galvanic), cathode = reduction (+ in galvanic). Notation: anode ‖ cathode.
- E°cell = E°cathode − E°anode; positive ⇒ spontaneous. SHE is the reference at 0.00 V.
- Nernst: E = E° − (0.059/n) log Q. At equilibrium E = 0 and Q = K.
- ΔG° = −nFE° (F = 96500); E° = (0.059/n) log K.
- κ = 1/ρ; Λm = κ × 1000/M. On dilution κ decreases, Λm increases.
- Strong electrolyte: Λm rises slowly (Λm = Λm° − A√C). Weak: rises steeply → use Kohlrausch for Λm°.
- α = Λm/Λm°; Ka = Cα²/(1 − α).
- Faraday: m = MIt/nF. Lead storage battery (rechargeable); H₂–O₂ fuel cell (only product water).
- Corrosion is electrochemical → prevent by galvanising and cathodic protection.
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