Chapter 2: Electrochemistry

Galvanic (Voltaic) Cell

• Converts Chemical Energy of a spontaneous redox reaction into Electrical Energy.
• Anode: Oxidation takes place here. It has a Negative (-) charge.
• Cathode: Reduction takes place here. It has a Positive (+) charge.
• Example: Daniell Cell (Zn-Cu cell).

Electrolytic Cell

• Converts Electrical Energy into Chemical Energy.
• Used to carry out non-spontaneous redox reactions by passing external current.
• Anode: Positive (+). Cathode: Negative (-).

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 the mechanical mixing of the two electrolytic solutions.

Standard 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.

E°(cell) = E°(cathode) - E°(anode)

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.
• Representation: Pt(s) | H2(g, 1 atm) | H+(aq, 1 M)

Electrochemical Series & Applications

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

• Strongest Oxidizing Agent: Fluorine (F2) - highest positive E° value.
• Strongest Reducing Agent: Lithium (Li) - highest negative E° value.
• Metals with negative E° value can displace Hydrogen gas from dilute acids.
• Helps in predicting the feasibility of a redox reaction (Reaction is feasible if E°(cell) is positive).

Nernst Equation

Relates electrode potential with the concentration of ions and temperature.

For a general reaction: aA + bB → cC + dD

E(cell) = E°(cell) - (2.303 RT / nF) log ([C]^c[D]^d / [A]^a[B]^b)

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

E(cell) = E°(cell) - (0.0591 / n) log Q

(where n = number of electrons transferred, Q = reaction quotient)

Equilibrium Constant (Kc)

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

E°(cell) = (0.0591 / n) log Kc

Gibbs Free Energy (ΔG)

Electrical work done in one second is equal to electrical potential multiplied by total charge passed.

ΔG = - n F E(cell) ΔG° = - n F E°(cell) = - 2.303 RT log Kc

(F = Faraday's constant = 96487 C ≈ 96500 C)

Resistance & Conductance

• Resistance (R): R = ρ(l/A). Unit: Ohm (Ω)
• Conductance (G): Reciprocal of resistance. G = 1/R = κ(A/l). Unit: Siemens (S) or Ohm⁻¹.
• Specific Conductivity (κ): κ = 1/ρ = G × (l/A). Unit: S cm⁻¹.
• Cell Constant (G*): Ratio of distance between electrodes and area of cross-section. G* = l/A.

Molar Conductivity (Λm)

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

Λm = (κ × 1000) / Molarity

Unit: S cm² mol⁻¹

Effect of Dilution: Both κ and Λm change with concentration. Specific conductivity (κ) decreases with dilution, while Molar conductivity (Λm) increases with dilution.

Kohlrausch's Law of Independent Migration of Ions

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

Λ°m = ν+λ°+ + ν-λ°-

Applications:

• Calculation of Λ°m for weak electrolytes.
• Calculation of degree of dissociation (α) = Λm / Λ°m.
• Calculation of dissociation constant (K) = (Cα²) / (1 - α).

First Law of Electrolysis

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

W ∝ Q ⇒ W = Z × Q W = Z × I × t

Where W = mass deposited, Z = electrochemical equivalent, I = current in Amperes, t = time in seconds.

Z = Equivalent Weight / 96500

Second Law of Electrolysis

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

W1 / W2 = E1 / E2

Primary Batteries (Non-Rechargeable)

• Dry Cell (Leclanche Cell): Anode is Zinc, Cathode is Graphite rod surrounded by MnO2 and Carbon. Electrolyte is paste of NH4Cl and ZnCl2. Voltage: 1.5 V.
• Mercury Cell: Used in hearing aids, watches. Anode is Zn-Hg amalgam. Voltage remains constant at 1.35 V throughout its life.

Secondary Batteries (Rechargeable)

• Lead Storage Battery: Used in automobiles and inverters.
  Anode: Spongy Lead (Pb). Cathode: Lead dioxide (PbO2). Electrolyte: 38% H2SO4.
  Discharging Reaction: Pb + PbO2 + 2H2SO4 → 2PbSO4 + 2H2O
• Nickel-Cadmium Cell: Longer life than lead storage battery but expensive.
  Reaction: Cd(s) + 2Ni(OH)3(s) → CdO(s) + 2Ni(OH)2(s) + H2O(l)

Fuel Cells

Galvanic cells designed to convert the energy of combustion of fuels like hydrogen, methane, methanol directly into electrical energy. Highly efficient and pollution-free.

• H2-O2 Fuel Cell Reaction: 2H2(g) + O2(g) → 2H2O(l)

Corrosion

Slow coating of surfaces of metallic objects with oxides or other salts. In rusting of iron, a miniature electrochemical cell is formed.

• Anode (Oxidation): 2Fe(s) → 2Fe²⁺ + 4e⁻
• Cathode (Reduction): O2(g) + 4H⁺(aq) + 4e⁻ → 2H2O(l)
• Prevention: Barrier protection (painting, greasing), Galvanization (coating with Zinc), and Sacrificial protection.