Chapter 3 of Class 12 Chemistry — Electrochemistry — connects chemistry with electricity. This chapter explains how chemical reactions can produce electric current (galvanic cells) and how electricity can drive non-spontaneous reactions (electrolytic cells). With 7-8 marks in Board exams, this is one of the highest-weightage chapters in Chemistry. Master the Nernst equation numericals and you’ll ace this!
Key Concepts
Electrochemical Cells
| Feature | Galvanic (Voltaic) Cell | Electrolytic Cell |
|---|---|---|
| Energy conversion | Chemical → Electrical | Electrical → Chemical |
| ΔG | Negative (spontaneous) | Positive (non-spontaneous) |
| E_cell | Positive | Negative (external EMF applied) |
| Anode | Negative terminal | Positive terminal |
| Cathode | Positive terminal | Negative terminal |
| Example | Daniel cell | Electrolysis of NaCl |
Daniel Cell — The Classic Galvanic Cell
Zn | Zn²⁺(aq) || Cu²⁺(aq) | Cu
- Anode (oxidation): Zn → Zn²⁺ + 2e⁻
- Cathode (reduction): Cu²⁺ + 2e⁻ → Cu
- Salt bridge: Maintains electrical neutrality (usually KCl in agar-agar)
- E⁰_cell = E⁰_cathode − E⁰_anode = +0.34 − (−0.76) = +1.10 V
Standard Electrode Potential (E⁰)
Measured against the Standard Hydrogen Electrode (SHE) which is assigned E⁰ = 0.00 V.
If E⁰_cell > 0 → reaction is spontaneous
If E⁰_cell < 0 → reaction is non-spontaneous
Electrochemical Series (selected values):
| Electrode | E⁰ (V) | Tendency |
|---|---|---|
| Li⁺/Li | −3.05 | Strongest reducing agent |
| K⁺/K | −2.93 | Strong reducing agent |
| Zn²⁺/Zn | −0.76 | Good reducing agent |
| H⁺/H₂ (SHE) | 0.00 | Reference |
| Cu²⁺/Cu | +0.34 | Weak oxidising agent |
| Ag⁺/Ag | +0.80 | Good oxidising agent |
| F₂/F⁻ | +2.87 | Strongest oxidising agent |
Nernst Equation
At 298 K: E_cell = E⁰_cell − (0.0591/n) log Q
Where: n = number of electrons transferred, Q = reaction quotient
At equilibrium: E_cell = 0, so E⁰_cell = (0.0591/n) log K_c
Relationship Between ΔG and E_cell
Where: n = moles of electrons, F = 96485 C/mol (Faraday constant)
If E⁰_cell > 0 → ΔG⁰ < 0 → spontaneous
If E⁰_cell < 0 → ΔG⁰ > 0 → non-spontaneous
Conductance of Electrolytic Solutions
Specific conductance (κ) = 1/ρ = G × (l/A) — unit: S/cm
Molar conductivity (Λ_m) = κ × 1000/c — unit: S·cm²/mol
Where c = concentration in mol/L
Variation of Conductivity with Concentration
- κ (specific conductance) decreases with dilution — fewer ions per unit volume
- Λ_m (molar conductivity) increases with dilution — more dissociation, less inter-ionic attraction
Kohlrausch’s Law of Independent Migration
Where ν = number of ions, λ⁰ = limiting molar conductivity of each ion
Example: Λ⁰_m(NaCl) = λ⁰(Na⁺) + λ⁰(Cl⁻)
Applications:
- Calculate Λ⁰_m of weak electrolytes (like CH₃COOH) which can’t be measured directly
- Λ⁰_m(CH₃COOH) = Λ⁰_m(CH₃COONa) + Λ⁰_m(HCl) − Λ⁰_m(NaCl)
- Determine degree of dissociation: α = Λ_m / Λ⁰_m
Electrolysis — Faraday’s Laws
Where: m = mass deposited, Z = electrochemical equivalent, I = current, t = time
Second Law: m₁/m₂ = E₁/E₂ (equivalent weights)
When same charge passes through different electrolytes
Batteries
| Battery | Type | Anode | Cathode | E_cell |
|---|---|---|---|---|
| Dry Cell (Leclanché) | Primary | Zn | MnO₂ + C | 1.5 V |
| Mercury Cell | Primary | Zn-Hg | HgO | 1.35 V |
| Lead-acid (car) | Secondary | Pb | PbO₂ | 2 V/cell |
| Ni-Cd | Secondary | Cd | NiO(OH) | 1.2 V |
| H₂-O₂ Fuel Cell | Fuel | H₂ | O₂ | 1.23 V |
Corrosion
Corrosion is an electrochemical process. Rusting of iron:
- Anode: Fe → Fe²⁺ + 2e⁻ (iron dissolves)
- Cathode: O₂ + 2H₂O + 4e⁻ → 4OH⁻
- Fe²⁺ + 2OH⁻ → Fe(OH)₂ → further oxidises to Fe₂O₃·xH₂O (rust)
Prevention: Galvanisation (Zn coating), electroplating, cathodic protection, painting, alloying
Important Definitions
| Term | Definition |
|---|---|
| Galvanic Cell | Device that converts chemical energy to electrical energy via spontaneous redox reaction |
| Electrolytic Cell | Device that uses electrical energy to drive non-spontaneous chemical reactions |
| Standard Electrode Potential | Potential of an electrode measured against SHE under standard conditions (1M, 1 atm, 25°C) |
| Salt Bridge | U-tube with electrolyte that maintains electrical neutrality in a galvanic cell |
| Molar Conductivity | Conductance of a solution containing 1 mole of electrolyte between electrodes 1 cm apart |
| Faraday Constant | Charge on 1 mole of electrons = 96485 C/mol |
| Corrosion | Electrochemical destruction of metals by reaction with the environment |
Solved Examples — NCERT Based
Example 1: EMF of Cell using Nernst Equation
Q: Calculate the EMF of the cell: Zn | Zn²⁺(0.001 M) || Cu²⁺(0.1 M) | Cu. Given E⁰(Zn²⁺/Zn) = −0.76 V, E⁰(Cu²⁺/Cu) = +0.34 V
Solution:
E⁰_cell = 0.34 − (−0.76) = 1.10 V, n = 2
Q = [Zn²⁺]/[Cu²⁺] = 0.001/0.1 = 0.01
E_cell = 1.10 − (0.0591/2) × log(0.01)
E_cell = 1.10 − 0.02955 × (−2) = 1.10 + 0.0591 = 1.159 V
Example 2: ΔG from E⁰_cell
Q: Calculate ΔG⁰ for the reaction: 2Fe³⁺ + 2I⁻ → 2Fe²⁺ + I₂. Given E⁰(Fe³⁺/Fe²⁺) = +0.77 V, E⁰(I₂/I⁻) = +0.54 V
Solution:
E⁰_cell = E⁰_cathode − E⁰_anode = 0.77 − 0.54 = 0.23 V
n = 2 (2 electrons transferred)
ΔG⁰ = −nFE⁰ = −2 × 96485 × 0.23 = −44,383 J = −44.4 kJ
Example 3: Faraday’s Law — Mass Deposited
Q: How much copper is deposited on the cathode if a current of 0.5 A is passed through CuSO₄ solution for 2 hours? (M of Cu = 63.5, n = 2)
Solution:
t = 2 × 3600 = 7200 s
m = (M × I × t) / (n × F) = (63.5 × 0.5 × 7200) / (2 × 96485)
m = 228600 / 192970 = 1.185 g
Example 4: Kohlrausch’s Law
Q: Calculate Λ⁰_m of CH₃COOH using: Λ⁰_m(CH₃COONa) = 91.0, Λ⁰_m(HCl) = 426.2, Λ⁰_m(NaCl) = 126.5 (all in S·cm²/mol)
Solution:
Λ⁰_m(CH₃COOH) = Λ⁰_m(CH₃COONa) + Λ⁰_m(HCl) − Λ⁰_m(NaCl)
= 91.0 + 426.2 − 126.5 = 390.7 S·cm²/mol
Important Questions for Board Exams
1 Mark Questions
- What is the SI unit of molar conductivity?
- How does molar conductivity change with dilution?
- What is the role of salt bridge in a galvanic cell?
- State Faraday’s first law of electrolysis.
- Why is E⁰ of SHE taken as zero?
2 Mark Questions
- Differentiate between galvanic and electrolytic cells.
- Write the Nernst equation and define each term.
- Explain why Λ_m of a weak electrolyte increases sharply at very low concentrations.
- State Kohlrausch’s law. Give one application.
3 Mark Questions
- Describe the construction and working of the Daniel cell with a diagram.
- How is ΔG related to EMF? Calculate ΔG⁰ for a cell with E⁰ = 1.10 V and n = 2.
- Explain corrosion of iron as an electrochemical process. How can it be prevented?
- Calculate the EMF of a cell at 298 K using the Nernst equation given specific conditions.
5 Mark Questions
- What is a fuel cell? Explain the construction and working of the hydrogen-oxygen fuel cell. What are its advantages?
- State and explain Faraday’s laws of electrolysis. Calculate the mass of aluminium deposited when 5 A current is passed through Al₂O₃ for 1 hour.
Quick Revision Points
- Anode = Oxidation, Cathode = Reduction (in ALL cells)
- E⁰_cell = E⁰_cathode − E⁰_anode; positive → spontaneous
- Nernst equation at 298K: E = E⁰ − (0.0591/n) log Q
- At equilibrium: E = 0, E⁰ = (0.0591/n) log K
- ΔG⁰ = −nFE⁰ (link between thermodynamics and electrochemistry)
- κ decreases with dilution; Λ_m increases with dilution
- Kohlrausch: Λ⁰_m = ν₊λ⁰₊ + ν₋λ⁰₋
- Faraday: m = MIt/nF
- Lead-acid: rechargeable, 6 cells × 2V = 12V car battery
- Fuel cells: high efficiency (~70%), clean energy, produce only water
- Corrosion = electrochemical; prevent by galvanisation, cathodic protection
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