Class 12 Chemistry Notes: Electrochemistry (Chapter 3)
Master Chapter 3 of Class 12 Chemistry. Understand galvanic vs electrolytic cells, Nernst equation, Faraday's laws, conductance, and electrolysis products with key formulas and exam PYQs.
Students know that electrochemistry involves electricity and chemical reactions. But they cannot consistently identify which electrode is the anode and which is the cathode, or predict what happens at each. This single confusion cascades into errors across every sub-topic.
1. The Universal Rule: Anode and Cathode
Do not memorise signs for anode and cathode separately for galvanic vs. electrolytic cells. Use the one rule that never changes:
The Universal Rule
Oxidation always occurs at the Anode.
Reduction always occurs at the Cathode.
| Property | Galvanic Cell | Electrolytic Cell |
|---|---|---|
| Anode sign | Negative (−) | Positive (+) |
| Cathode sign | Positive (+) | Negative (−) |
| Energy conversion | Chemical → Electrical | Electrical → Chemical |
2. Cell EMF
E°cell = E°cathode − E°anode
Always subtract anode from cathode. If E°cell is negative, the reaction is non-spontaneous as written. Students who add instead of subtract, or subtract in the wrong order, arrive at wrong spontaneity conclusions.
3. Nernst Equation
Formula: Nernst Equation (at 298 K)
E = E° − (0.0592/n) log Q
- n = number of electrons transferred in the balanced equation
- Q = reaction quotient (products/reactants, excluding pure solids and liquids)
Common errors: using log instead of ln in the full form, using the wrong n, or writing Q incorrectly by including solids.
4. Faraday's Laws of Electrolysis
The step-by-step calculation method:
- Calculate charge: Q = I × t (coulombs)
- Calculate moles of electrons: mol e⁻ = Q / 96485
- Calculate moles of substance deposited: mol substance = mol e⁻ / n (where n = electrons per ion)
- Calculate mass: mass = moles × molar mass
⚠️ Watch Out! — The Value of n
For Cu²⁺, n = 2. Depositing 1 mole of copper requires 2 Faradays of charge. Students who use n = 1 for all metals get answers that are exactly double or half the correct value.
Always check the ion charge before calculating.
5. The Salt Bridge
Its function is often stated but not understood. As a galvanic cell operates:
- Positive ions accumulate near the cathode → the solution would become positively charged.
- Negative ions accumulate near the anode → the solution would become negatively charged.
The salt bridge provides ions to neutralise these charges, allowing the cell to continue operating. Without a salt bridge, the cell voltage drops to zero almost immediately.
Summary: Formula Sheet
| Concept | Formula |
|---|---|
| Cell EMF | E°cell = E°cathode − E°anode |
| Nernst equation (298 K) | E = E° − (0.0592/n) log Q |
| Charge | Q = I × t |
| Faraday's constant | F = 96485 C/mol |
| Conductivity (dilution) | Decreases (fewer ions per unit volume) |
| Molar conductivity (dilution) | Increases (ions more mobile, weak electrolytes dissociate more) |
Practice Questions (PYQs)
- Define standard electrode potential. How is the standard cell EMF calculated from electrode potentials?
- A current of 2 A is passed through a solution of CuSO₄ for 30 minutes. Calculate the mass of copper deposited at the cathode. (Atomic mass of Cu = 63.5 g/mol, F = 96485 C/mol)
- Write the Nernst equation for the cell: Zn | Zn²⁺ (0.1 M) || Cu²⁺ (0.01 M) | Cu. Calculate E at 298 K if E°cell = 1.10 V.
- Why does the molar conductivity of a weak electrolyte increase sharply on dilution, while that of a strong electrolyte increases only slightly?
- Explain the function of a salt bridge in a galvanic cell. What would happen if you replaced it with a piece of glass rod?