Electrochemistry

Electrochemistry Basics

Electrochemistry deals with the relationship between electrical energy and chemical energy changes in redox reactions.

Electrical Resistance & Conductance

1) Resistance (R)

• Resistance measures obstruction to the flow of current.
• R ∝ l/a → R = ρ(l/a)
• ρ = resistivity (specific resistance)
• Unit of resistance: ohm (Ω)

2) Resistivity / Specific Resistance (ρ)

• ρ = R(a/l)
• If l = 1 cm and a = 1 cm², then R = ρ
• Unit (cgs): Ω·cm; Unit (SI): Ω·m

3) Conductance (C)

• Conductance measures ease of current flow; it is additive.
• C = 1/R
• Unit (cgs): mho or Ω⁻¹; Unit (SI): Siemens (S)

4) Conductivity / Specific Conductance (κ)

• Conductivity is inverse of resistivity: κ = 1/ρ
• Defined as conductance of a solution of length 1 cm and area 1 cm².
• Unit (cgs): Ω⁻¹ cm⁻¹; Unit (SI): S m⁻¹

5) Molar Conductivity / Molar Conductance (Λ_m)

• Conducting power of all ions produced by dissolving 1 mole of electrolyte.
• Λ_m = κ × 1000 / M
• Unit (cgs): Ω⁻¹ cm² mol⁻¹; Unit (SI): S m² mol⁻¹

6) Equivalent Conductivity (Λ_eq)

• Conducting power of all ions produced by dissolving 1 gram equivalent of electrolyte.
• Λ_eq = κ × 1000 / N (N = normality)
• Unit (cgs): Ω⁻¹ cm² (g equiv)⁻¹; Unit (SI): S m² equiv⁻¹

7) Conductivity from Cell Constant

• Conductivity = Conductance × Cell constant
• Cell constant = l/a
• Unit: cm⁻¹ (cgs) or m⁻¹ (SI)

Factors Affecting Electrolytic Conductance

• Nature of electrolyte
• Concentration: molar conductance increases on dilution (decrease in concentration)
• Temperature: conductivity increases with increase in temperature

Kohlrausch’s Law

Kohlrausch’s law: Equivalent conductivity at infinite dilution is the sum of the equivalent conductivities of cations and anions.

• Λ⁰(A_xB_y) = x λ⁰_A⁺ + y λ⁰_B⁻

Applications (names)

• Determination of Λ⁰ for weak electrolytes
• Degree of ionisation of weak electrolyte
• Ionisation constant of weak electrolyte
• Solubility of sparingly soluble salts

Transport (Transference) Number

Transport number = fraction of total current carried by an ion.

• t_anion = current carried by anion / total current
• t_cation = current carried by cation / total current
• t_anion + t_cation = 1

Determination (names)

• Hittorf’s method
• Moving boundary method
• EMF method
• From ionic mobility

Types of Cells (Quick Compare)

• Primary cells: cannot be recharged (dry cell, mercury cell)
• Secondary cells: rechargeable (lead storage battery, Ni–Cd cell)
• Fuel cells: energy from combustion converted to electrical energy (H₂–O₂ fuel cell)

Main Features of Common Cells

Conductors, Electrolytes & Electrolysis

Conductors vs Insulators

• Conductors allow current to pass; insulators do not.

Electronic (Metallic) vs Electrolytic Conductors

Key Definitions

• Electrolyte: conducts electricity in molten/solution state and is decomposed by it.
• Electrode (half-cell): electronic conductor in contact with an electrolytic conductor.
• Electrode potential: potential difference at interface of electronic and electrolytic conductors.
• Cell: assembly of two half-cells.

Electrolysis (Core Rules)

• Cations migrate to cathode and are reduced: Mⁿ⁺ + ne⁻ → M
• Anions migrate to anode and are oxidised: Aⁿ⁻ → A + ne⁻
• Products depend on:
  • Nature of electrodes (attackable / non-attackable)
  • Nature of electrolyte (molten / aqueous)
  • Charge density passed
  • Concentration of solution

Examples: Discharge Products (as given)

Electrolysis Notes

• Electrolysis occurs only at electrodes: oxidation at anode, reduction at cathode.
• If two or more similar ions compete for discharge: higher discharge potential → lower tendency to discharge.
• Discharge occurs only during passage of charge.

Faraday’s Laws of Electrolysis

First Law

• Mass deposited (w) is proportional to charge (Q): w ∝ Q, where Q = it
• w = Z × i × t (Z = electrochemical equivalent)

Second Law

• For same charge passed through different electrolytes: w ∝ E (E = equivalent weight)
• w₁/w₂ = E₁/E₂ (Q constant)

Quick Notes

• Passing 1 coulomb deposits Z g of substance.
• Passing 1 faraday deposits E g of substance.

Applications of Electrolysis (names)

• Extraction of metals
• Preparation of chemicals
• Preparation of organic compounds
• Corrosion and its prevention
• Purification of metals

Electrolytic Cells vs Electrochemical Cells

Electrochemical Cells

Chemical Cells (Daniel Cell)

• Representation: Zn | Zn²⁺ || Cu²⁺ | Cu
• Anode (oxidation): Zn → Zn²⁺ + 2e⁻
• Cathode (reduction): Cu²⁺ + 2e⁻ → Cu
• Overall: Zn + Cu²⁺ → Zn²⁺ + Cu
• Electrons released at anode flow through external circuit to cathode.

Concentration Cells

• No net chemical change; cell works due to concentration difference.
• General form: M | M⁺(C1) || M⁺(C2) | M where C2 > C1
• Electrode in more dilute solution acts as anode.

Nernst Equation (at 25°C)

For Metal Electrode

• General: M ⇌ Mⁿ⁺ + ne⁻
• E = E° − (0.0591/n) log(a_OS/a_RS)

For Non-metal / Anion Electrode

• General: A + ne⁻ ⇌ Aⁿ⁻
• E = E° − (0.0591/n) log(a_RS/a_OS)

Applications of Nernst Equation

(i) EMF of cell

• E°_cell = E°(OP) + E°(RP)
• E_cell = E(OP) + E(RP)

(ii) EMF and Equilibrium Constant

• ΔG° = −nFE°
• 0.0591 logK / n = E° (at 25°C)

(iii) Heat of reaction (as given)

• ΔH = nF [ T(∂E/∂T)_P − E ]
• (∂E/∂T)_P = temperature coefficient of EMF

Other uses (names)

• To decide spontaneity of cell reaction
• To evaluate solubility product
• To evaluate pH

Reversible vs Irreversible Cells (Test)

• If E(test) slightly > E(external): current flows from test to external; reaction occurs.
• If E(test) slightly < E(external): current flows external to test; reaction reverses.
• If E(test) = E(external): no current; no reaction.

Electrochemical Series (Key Points)

• Arranged by tendency to lose electrons (oxidation): decreasing trend of E°(OP)
• More +ve E°(OP) → higher tendency to get oxidised
• Stronger oxidant has weaker conjugate reductant (and vice versa)
• Reducing power of metals decreases down the series
• Oxidising power of metal ions increases down the series
• Metals above hydrogen displace H₂ from dilute acids
• Metal higher in series displaces metal below from its salt solution
• Halide reducing power: HI > HBr > HCl > HF

Lead Storage Battery & Corrosion (Quick Notes)

Lead Storage Battery (Discharge)

• Anode: Pb + SO₄²⁻ → PbSO₄ + 2e⁻
• Cathode: PbO₂ + SO₄²⁻ + 4H⁺ + 2e⁻ → PbSO₄ + 2H₂O
• Overall: Pb + PbO₂ + 2H₂SO₄ → 2PbSO₄ + 2H₂O
• As H₂SO₄ is consumed, battery voltage drops.
• Charging reactions are reverse of discharge.

Corrosion

• Metal surface changes to oxides/sulphides/carbonates due to atmospheric attack.
• Enhanced by: impurities, moisture, electrolytes (e.g., saline water).
• Prevention:
  • Barrier protection: oil/grease, paints, electroplating
  • Sacrificial protection: coating with more electropositive metal (e.g., Zn) — galvanisation
  • Electrical protection: connect iron to more electropositive metal using a wire

Redox Quick Revision

• Loss of electron = oxidation; gain of electron = reduction
• Increase in oxidation number = oxidation; decrease = reduction
• Oxidised substance acts as reducing agent; reduced substance acts as oxidising agent

Equivalent Mass

• Equivalent mass of oxidising agent = (Formula mass of O.A.) / (No. of electrons gained)
• Equivalent mass of reducing agent = (Formula mass of R.A.) / (No. of electrons lost)

Types of Redox Reactions (examples)

• Intermolecular: Zn + 2HCl → ZnCl₂ + H₂
• Intramolecular: 2KClO₃ → 2KCl + 3O₂
• Disproportionation: 2Cu⁺ → Cu²⁺ + Cu