The direct answer
Galvanic vs Electrolytic Cells
| Feature | Galvanic (Voltaic) | Electrolytic |
|---|---|---|
| Energy conversion | Chemical → Electrical | Electrical → Chemical |
| Spontaneity | Spontaneous (ΔG < 0) | Non-spontaneous (ΔG > 0) |
| Anode | Negative (−), oxidation | Positive (+), oxidation |
| Cathode | Positive (+), reduction | Negative (−), reduction |
| Example | Daniel cell, batteries | Electroplating, electrolysis |
🧠 Memory trick: "An Ox, Red Cat" — Anode is Oxidation, Reduction at Cathode. In galvanic cells: anode is negative. In electrolytic: anode is positive.
The Nernst Equation
General form: E = E° − (RT/nF)·ln Q
At 298K (25°C): E = E° − (0.0591/n)·log Q
For a cell reaction: aA + bB → cC + dD
Q = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ
Key relationships:
ΔG = −nFE (spontaneous when E > 0)
ΔG° = −nFE°
Equilibrium: E = 0 at equilibrium, E° = (0.0591/n)·log K
🧠 Memory trick: "E = E° − 0.0591/n log Q" — at 298K. This single formula solves most JEE electrochemistry numericals. At equilibrium, E = 0 and log K = nE°/0.0591.
Faraday's Laws of Electrolysis
First Law: m = Z·Q = Z·I·t
Mass deposited = electrochemical equivalent × charge
Second Law: m₁/m₂ = E₁/E₂ (equivalent weights)
Key values:
1 Faraday (F) = 96,500 C/mol e⁻
Z = E/96,500 (gram equivalent per coulomb)
Charge = I × t (Coulombs)
Example: For Cu²⁺ + 2e⁻ → Cu, 2F deposits 63.5g of Cu. 1F deposits 31.75g.
🧠 Memory trick: "m = ZIt" — mass equals electrochemical equivalent times current times time. One Faraday = 96,500 coulombs = one mole of electrons.
Conductance and Molar Conductivity
Key definitions:
Conductance (G) = 1/R (siemens, S)
Conductivity (κ) = G × (l/a) (S/m)
Molar conductivity: Λₘ = κ/c (c in mol/m³)
Λₘ = κ × 1000/M (c in mol/L, κ in S/cm)
Kohlrausch's law: Λₘ = Λₘ° − A√c (strong electrolytes)
Λₘ° values: H⁺ = 349.8, OH⁻ = 198.5, Na⁺ = 50.1 (S cm²/mol)
🧠 Memory trick: "Conductivity decreases with concentration, but molar conductivity increases." For strong electrolytes, Λₘ vs √c is linear (Kohlrausch).
Complete Electrochemistry Formula Sheet
| Quantity | Formula | Units |
|---|---|---|
| EMF of cell | E°(cell) = E°(cathode) − E°(anode) | Volts (V) |
| Nernst equation | E = E° − (0.0591/n)·log Q | Volts (V) |
| Gibbs free energy | ΔG = −nFE | Joules (J) |
| Equilibrium constant | log K = nE°/0.0591 | — |
| Faraday's 1st law | m = Z·I·t | grams (g) |
| Charge | Q = I·t | Coulombs (C) |
| Molar conductivity | Λₘ = κ × 1000/M | S cm²/mol |
| Kohlrausch | Λₘ = Λₘ° − A√c | S cm²/mol |
| 1 Faraday | F = 96,500 C | C/mol e⁻ |
JEE PYQ Patterns on Electrochemistry
| Question Type | Frequency | How to Approach |
|---|---|---|
| Nernst equation calculations | Very High | Identify n (electrons transferred), write Q, apply E = E° − (0.0591/n)log Q |
| Faraday's law mass deposition | Very High | m = ZIt, use equivalent weights for multi-electrode cells |
| EMF from electrode potentials | High | E°(cell) = E°(cathode) − E°(anode), identify which is which |
| Conductance problems | Medium | Λₘ = κ × 1000/M, Kohlrausch's law for limiting values |
| Spontaneity | Medium | E > 0 → spontaneous, ΔG < 0 → spontaneous |
| Batteries and fuel cells | Medium | Know reactions of Daniel cell, lead-acid, fuel cells |
Common JEE traps to avoid:
❌ Using wrong n value in Nernst equation (count electrons actually transferred)
❌ Forgetting to balance the cell reaction before using Q
❌ Mixing up anode/cathode signs in galvanic vs electrolytic cells
❌ Not converting 0.0591 vs 0.059 (it's 0.0591 at 298K)
❌ Using concentration in mol/L vs mol/m³ for conductivity
Frequently Asked Questions
What is the Nernst equation?
The Nernst equation relates cell potential to concentration: E = E° − (0.0591/n)·log Q at 298K. Where E° is standard cell potential, n is number of electrons transferred, and Q is the reaction quotient. This is essential for JEE electrochemistry problems.
What is Faraday's first law of electrolysis?
Faraday's first law states that the mass of substance deposited at an electrode is proportional to the quantity of electricity passed: m = Z·Q = Z·I·t. Where Z is electrochemical equivalent (Z = E/96500), I is current in amperes, and t is time in seconds.
What is the difference between galvanic and electrolytic cells?
Galvanic (voltaic) cells convert chemical energy to electrical energy spontaneously (ΔG < 0). Electrolytic cells use electrical energy to drive non-spontaneous reactions (ΔG > 0). Galvanic cells have anode negative, cathode positive. Electrolytic cells have anode positive, cathode negative.
What is standard electrode potential?
Standard electrode potential (E°) is the potential of an electrode measured against the standard hydrogen electrode (SHE, E° = 0V) under standard conditions (1M, 1 bar, 25°C). Higher E° means greater tendency to be reduced. More positive E° = stronger oxidizing agent.
How do I calculate EMF of a cell?
EMF = E°(cathode) − E°(anode). The cathode is where reduction occurs (higher reduction potential). The anode is where oxidation occurs (lower reduction potential). For non-standard conditions, use the Nernst equation: E = E° − (0.0591/n)·log Q.
What is conductance and molar conductivity?
Conductance (G) = 1/R = κ·(a/l), where κ is conductivity. Molar conductivity Λₘ = κ/c, where c is concentration in mol/m³. For strong electrolytes, Λₘ = Λₘ° − A√c (Kohlrausch's law). Λₘ increases as concentration decreases.
What are the most common electrochemistry questions in JEE?
JEE frequently asks: (1) Nernst equation calculations, (2) Faraday's laws for mass deposition, (3) EMF calculation from electrode potentials, (4) Predicting spontaneity from ΔG, (5) Conductivity and molar conductivity, (6) Electrolysis products prediction, (7) Battery and fuel cell questions.
Explore more chemistry resources
Master electrochemistry, then move to chemical kinetics and thermodynamics.
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