Electrochemistry

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your exam.

Electrochemistry — Quick Facts

Electrochemistry deals with the interconversion of chemical energy and electrical energy. Galvanic cells convert chemical energy to electricity; electrolytic cells use electricity to drive non-spontaneous reactions.

Key Definitions:

• Anode: Electrode where oxidation occurs (loss of electrons)
• Cathode: Electrode where reduction occurs (gain of electrons)
• Voltaic/Galvanic cell: Spontaneous reaction → electricity generated
• Electrolytic cell: Non-spontaneous reaction → electricity consumed

Essential Formulas:

• Cell emf: E°_cell = E°_cathode − E°_anode
• Nernst equation: E = E° − (0.0591/n) log Q at 298 K
• For the reaction: aA + bB → cC + dD, Q = [C]^c[D]^d / [A]^a[B]^b
• Free energy: ΔG = −nFE (F = 96500 C mol⁻¹)
• Equilibrium constant: ΔG° = −RT ln K → E°_cell = (RT/nF) ln K = (0.0591/n) log K

⚡ Exam tip: When writing Nernst equation, always use n = total electrons transferred (not per species). For Zn + Cu²⁺ → Zn²⁺ + Cu, n = 2.

⚡ Common mistake: Students forget the sign convention. E°_cell must be positive for a spontaneous galvanic cell. If E°_cathode < E°_anode, the cell won’t work as written.

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🟡 Standard — Regular Study (2d–2mo)

Standard content for students with a few days to months.

Electrochemistry — NEET/JEE Study Guide

Standard Hydrogen Electrode (SHE): The SHE is assigned E° = 0.00 V at 1 M H⁺ and 1 atm H₂ gas. All other standard reduction potentials are measured relative to it.

Galvanic vs Electrolytic Cells:

Feature	Galvanic Cell	Electrolytic Cell
Energy	Chemical → Electrical	Electrical → Chemical
ΔG	Negative (spontaneous)	Positive (non-spontaneous)
Anode	Negative electrode	Positive electrode
Cathode	Positive electrode	Negative electrode
E_cell	Positive	Requires external source

Conductivity and Resistivity:

• κ (conductivity) = 1/ρ, unit S cm⁻¹
• Molar conductivity: Λ_m = κ × V_m (where V_m is volume containing 1 mole of electrolyte)
• Kohlrausch’s law: Λ_m° = ν₊λ₊° + ν₋λ₋° (sum of ionic conductivities)
• For weak electrolytes: Λ_m = Λ_m° × α, where α is degree of dissociation

Electrochemical Series (Selected Standard Potentials at 25°C):

Half-reaction	E° (V)
F₂ + 2e⁻ → 2F⁻	+2.87
O₂ + 4H⁺ + 4e⁻ → 2H₂O	+1.23
Ag⁺ + e⁻ → Ag	+0.80
Fe³⁺ + e⁻ → Fe²⁺	+0.77
O₂ + 2H₂O + 4e⁻ → 4OH⁻	+0.40
Cu²⁺ + 2e⁻ → Cu	+0.34
2H⁺ + 2e⁻ → H₂	0.00
Pb²⁺ + 2e⁻ → Pb	−0.13
Sn²⁺ + 2e⁻ → Sn	−0.14
Ni²⁺ + 2e⁻ → Ni	−0.25
Fe²⁺ + 2e⁻ → Fe	−0.44
Zn²⁺ + 2e⁻ → Zn	−0.76
Al³⁺ + 3e⁻ → Al	−1.66
Mg²⁺ + 2e⁻ → Mg	−2.37
Na⁺ + e⁻ → Na	−2.71
Ca²⁺ + 2e⁻ → Ca	−2.87

Faraday’s Laws of Electrolysis:

• 1st law: Mass deposited m = (Q × M) / (n × F) where Q = I × t
• 2nd law: Same Q through different electrolytes → masses proportional to equivalent masses
• 1 Faraday = 96500 C = 1 mole of electrons

⚡ NEET 2021 Qn: What mass of Ag is deposited when 1 A current flows for 16 minutes 40 seconds through AgNO₃ solution? m = (I × t × M) / (n × F) = (1 × 1000 × 108) / (1 × 96500) = 1.119 g

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🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for students on a longer study timeline.

Electrochemistry — Comprehensive Notes

Derivation of Nernst Equation: For a general redox reaction at equilibrium: ΔG = ΔG° + RT ln Q Since ΔG = −nFE and ΔG° = −nFE°: −nFE = −nFE° + RT ln Q E = E° − (RT/nF) ln Q

At 298 K: E = E° − (0.0591/n) log₁₀ Q (using log₁₀ and converting ln to log₁₀)

Concentration Cell: When both electrodes are the same metal in solutions of different concentrations. E°_cell = 0 for identical half-cells. E_cell depends only on the concentration ratio: E_cell = (0.0591/n) log([reducing agent]ₐ / [reducing agent]ᵦ)

This is also the basis of the glass electrode used in pH measurement.

Batteries:

1. Primary battery (non-rechargeable): LeclanchÉ cell (dry cell) — Zn anode, MnO₂ cathode, NH₄Cl electrolyte. E ≈ 1.5 V. Can’t be recharged.

2. Secondary battery (rechargeable):

   • Lead-acid battery: Pb anode, PbO₂ cathode, H₂SO₄ electrolyte. Each cell = 2 V. E_cell° = 1.85 V.
   • Reactions: Anode: Pb + SO₄²⁻ → PbSO₄ + 2e⁻ Cathode: PbO₂ + 4H⁺ + SO₄²⁻ + 2e⁻ → PbSO₄ + 2H₂O

3. Fuel cells: O₂ + H₂ → H₂O with Pt electrodes. E° = 1.23 V. High efficiency (~70%), used in spacecraft.

Corrosion: Corrosion is the electrochemical oxidation of metals. Rusting of iron requires both oxygen and water: Anode: Fe → Fe²⁺ + 2e⁻ Cathode: O₂ + 4H⁺ + 4e⁻ → 2H₂O (in acidic conditions) or: O₂ + 2H₂O + 4e⁻ → 4OH⁻ (in neutral/alkaline conditions) Prevention: Galvanising (coating with Zn), painting, or sacrificial anodes.

pH Determination using Hydrogen Electrode: Using Nernst for 2H⁺ + 2e⁻ → H₂: E = 0 − (0.0591/1) log(1/[H⁺]²) = 0 + 0.0591 log[H⁺] = −0.0591 pH This is the principle of the glass electrode in pH meters.

NEET Pattern Analysis: Electrochemistry contributes approximately 2 questions per year in NEET (4 marks). Focus areas: Nernst equation calculations, EMF prediction from standard potentials, Faraday’s laws, and identification of galvanic vs electrolytic cells. The product of electrolysis questions (what gets deposited at cathode/anode) are frequently tested when multiple ions are present — use reduction potentials to determine which species reduces/oxidises first.

⚡ NEET 2023 Qn: Calculate E_cell for Fe(s) | Fe²⁺(0.001 M) || Cu²⁺(0.1 M) | Cu(s). E°_cell = 0.34 − (−0.44) = 0.78 V n = 2 E_cell = 0.78 − (0.0591/2) log([Fe²⁺]/[Cu²⁺]) E_cell = 0.78 − 0.02955 log(0.001/0.1) = 0.78 − 0.02955 × (−2) = 0.78 + 0.0591 = 0.839 V