Electrochemistry

1. Introduction

Electrochemistry deals with the interconversion of electrical and chemical energy. It encompasses galvanic cells, electrolysis, conductance, and industrial applications.

2. Electrochemical Cells

2.1 Galvanic (Voltaic) Cells

Spontaneous reaction produces electricity. The cell consists of two half-cells connected by a salt bridge.

Anode: Oxidation occurs (negative terminal). Cathode: Reduction occurs (positive terminal).

2.2 Cell Notation

Zn | Zn²⁺ || Cu²⁺ | Cu. Single vertical line: phase boundary. Double line: salt bridge.

2.3 Standard Electrode Potential

The potential of a half-cell under standard conditions (1 M, 1 bar, 298 K) relative to the standard hydrogen electrode (SHE = 0 V).

E_cell⁰ = E_cathode⁰ - E_anode⁰

3. Nernst Equation

E_cell = E_cell⁰ - (0.0591/n) log Q (at 298 K)

For a cell: E = E⁰ - (RT/nF) ln Q

4. Electrolysis

Electrolysis is the decomposition of an electrolyte by passing electric current through it.

4.1 Faraday's Laws

First Law: The mass of substance deposited at an electrode is directly proportional to the quantity of electricity passed. m = ZIt, where Z = electrochemical equivalent.

Second Law: When the same quantity of electricity is passed through different electrolytes, the masses deposited are proportional to their chemical equivalents.

m₁/m₂ = E₁/E₂, where E = equivalent mass.

4.2 Applications of Electrolysis

• Electrorefining of metals (copper purification).
• Electroplating (coating one metal with another).
• Production of aluminium (Hall-Heroult process).
• Production of chlorine and NaOH (chlor-alkali industry).

5. Conductance

4.1 Types

Metallic conductance: Flow of electrons. Decreases with temperature. Electrolytic conductance: Flow of ions. Increases with temperature.

4.2 Molar Conductivity

Λ_m = κ/C, where κ is conductivity and C is concentration.

Λ_m increases with dilution because ions experience less interionic attraction.

4.3 Kohlrausch's Law

Λ_m⁰ = ν₊λ₊⁰ + ν₋λ₋⁰, where λ⁰ are limiting molar conductivities of individual ions.

Used to calculate Λ_m⁰ for weak electrolytes and degree of dissociation.

5. Batteries

5.1 Primary Cells (Non-rechargeable)

Leclanche cell: Zn (anode), NH₄Cl paste, MnO₂/C (cathode). EMF ≈ 1.5 V.

5.2 Secondary Cells (Rechargeable)

Lead storage battery: Pb (anode), PbO₂ (cathode), H₂SO₄ electrolyte. EMF ≈ 2 V per cell.

5.3 Fuel Cells

Generate electricity from fuel (H₂) and oxidant (O₂). H₂-O₂ fuel cell produces water as byproduct. Efficiency > 60%.

6. Corrosion

The gradual destruction of metals by electrochemical reactions. Rusting of iron requires oxygen and water.

Prevention: Alloying, galvanization, electroplating, painting, cathodic protection.

7. Worked Problems

Problem 1: Calculate E⁰ for Zn|Zn²⁺||Cu²⁺|Cu given E⁰_Zn²⁺/Zn = -0.76V and E⁰_Cu²⁺/Cu = 0.34V. Solution: E_cell⁰ = E_cathode⁰ - E_anode⁰ = 0.34 - (-0.76) = 1.10 V.

Problem 2: Calculate the EMF of the cell at 298 K: Zn|Zn²⁺(0.1M)||Cu²⁺(0.01M)|Cu. E_cell⁰ = 1.10V. Solution: E = 1.10 - (0.0591/2) log(0.1/0.01) = 1.10 - 0.0296 log(10) = 1.10 - 0.0296 = 1.07 V.

8. Common Mistakes

'Students often reverse the Nernst equation expression for Q. For Zn + Cu²⁺ → Zn²⁺ + Cu, Q = [Zn²⁺]/[Cu²⁺], not the reverse.'

9. ISC Exam Focus

10. Self-Test Questions

1. How does the molar conductivity of a strong electrolyte vary with dilution? Explain.
2. Calculate the EMF of the cell: Fe|Fe²⁺(0.1M)||Ag⁺(0.01M)|Ag. Given E⁰_Fe²⁺/Fe = -0.44V, E⁰_Ag⁺/Ag = 0.80V.
3. State and explain Kohlrausch's law of independent migration of ions.
4. Distinguish between primary and secondary batteries with examples.
5. Explain the mechanism of corrosion of iron. How is it prevented?