Chapter 2: Electrochemistry

Galvanic (Voltaic) Cell

• Converts Chemical Energy of a spontaneous redox reaction into Electrical Energy.
• Anode (Left): Oxidation takes place here. It has a Negative (-) charge.
• Cathode (Right): Reduction takes place here. It has a Positive (+) charge.
• Representation: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

Salt Bridge & Its Functions

A U-shaped tube containing an inert electrolyte (like KCl, KNO₃, or NH₄NO₃) in agar-agar gel.

• Completes the inner electrical circuit.
• Maintains the electrical neutrality of the two half-cells.
• Prevents liquid-liquid junction potential.

Electrode Potential (E°)

The potential difference developed between metal electrode and its ions in solution at 1 M concentration, 1 atm pressure, and 298 K temperature.

Standard Hydrogen Electrode (SHE)

• Used as a reference electrode. Its standard electrode potential is arbitrarily taken as Zero Volts (0.00 V).
• Consists of a platinum wire coated with platinum black, dipped in 1M H⁺ solution, with pure H₂ gas bubbled at 1 atm.

Electrochemical Series

Arrangement of elements in increasing order of their standard reduction potentials.

• Strongest Oxidizing Agent: Fluorine (F₂) - highest positive E° value (+2.87 V).
• Strongest Reducing Agent: Lithium (Li) - highest negative E° value (-3.05 V).
• Metals with negative E° value can displace Hydrogen gas from dilute acids.

Nernst Equation

Relates electrode potential with the concentration of ions and temperature. For a general reaction: aA + bB → cC + dD

At 298 K (25°C), the equation simplifies to:

• Concentration Cell: Both electrodes are of the same material but different concentrations. Here, E°(cell) = 0.

Equilibrium Constant (Kc) & Gibbs Energy

At equilibrium, the cell potential becomes zero (E(cell) = 0).

Note: If ΔG° is negative, E°(cell) is positive, and the reaction is spontaneous. (F = 96500 C)

Resistance, Conductance & Cell Constant

• Conductance (G): Reciprocal of resistance. G = 1/R = κ(A/l). Unit: Siemens (S).
• Specific Conductivity (κ): κ = 1/ρ = G × (l/A). Unit: S cm⁻¹. Decreases with dilution.
• Cell Constant (G*): Ratio of distance between electrodes to area. G* = l/A.

Molar Conductivity (Λm)

Conducting power of all the ions produced by dissolving one mole of an electrolyte. Increases with dilution.

Kohlrausch's Law

Limiting molar conductivity of an electrolyte can be represented as the sum of the individual contributions of the anion and cation.

• Applications: Calculate Λ°m for weak electrolytes (like CH₃COOH).
• Degree of dissociation: α = Λm / Λ°m
• Dissociation constant: Ka = (Cα²) / (1 - α)

First Law

The mass of any substance deposited at any electrode is directly proportional to the quantity of electricity passed.

Where W = mass, Z = electrochemical equivalent (Equivalent Weight / 96500), I = current (A), t = time (s).

Second Law

When the same quantity of electricity is passed through different electrolytes connected in series, the masses deposited are proportional to their equivalent weights.

Primary Batteries (Non-Rechargeable)

• Dry Cell: Anode is Zinc container, Cathode is Graphite rod. Electrolyte is paste of NH₄Cl + ZnCl₂. Voltage: 1.5 V.
• Mercury Cell: Used in hearing aids. Voltage remains constant at 1.35 V because overall reaction does not involve any ion in solution whose concentration can change.

Secondary Batteries (Rechargeable)

Lead Storage Battery: Used in cars. Anode = Spongy Lead. Cathode = PbO₂. Electrolyte = 38% H₂SO₄.

Fuel Cells

Convert combustion energy of fuels (H₂, CH₄) directly to electricity. Pollution-free. Efficiency is ~70%.

Corrosion (Rusting of Iron)

A miniature electrochemical cell forms on the iron surface.

• Anode: 2Fe(s) → 2Fe²⁺ + 4e⁻
• Cathode: O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l)
• Fe²⁺ is further oxidized by atmosphere to Fe³⁺ forming Rust: Fe₂O₃·xH₂O.