Statement A — TRUE ✓
In a galvanic cell (spontaneous), oxidation occurs at the anode (releases electrons → negative electrode) and reduction at the cathode (gains electrons → positive electrode). External circuit: electrons flow from anode (−) to cathode (+).
Statement B — FALSE ✗
In an electrolytic cell, the anode is connected to the POSITIVE terminal of the battery, and the cathode to the NEGATIVE terminal. Anode = positive, cathode = negative. This is opposite to galvanic cell. Oxidation still at anode, reduction still at cathode — but the electrode signs are reversed.
Statement C — TRUE ✓
The Standard Hydrogen Electrode (SHE) is the universal reference with E° = 0.00 V by convention. All other standard electrode potentials are measured relative to SHE. Conditions: H₂ gas at 1 atm, H⁺ at 1 M concentration, 25°C.
Statement D — TRUE ✓
E°cell = E°cathode − E°anode (reduction potentials). This is the standard formula for cell EMF. Since cathode has higher reduction potential than anode for a spontaneous cell, E°cell is positive.
📌 Galvanic cell: converts chemical energy → electrical energy (spontaneous)
📌 Electrolytic cell: converts electrical energy → chemical energy (non-spontaneous)
📌 Both: oxidation at anode, reduction at cathode
📌 Galvanic: anode(−), cathode(+) | Electrolytic: anode(+), cathode(−)
📌 Daniel cell: Zn|ZnSO₄||CuSO₄|Cu — classic galvanic cell
E°cell = E°cathode − E°anode
ΔG° = −nFE°cell
E°cell = (RT/nF) ln K = (0.0591/n) log K at 25°C
A positive E°cell means the reaction is spontaneous (ΔG° < 0). The more positive E°cell, the more spontaneous. The standard hydrogen electrode (SHE) has E° = 0.00 V by definition — all other E° values are measured relative to it.
The Nernst equation relates cell EMF to concentration of reactants and products:
E = E° − (RT/nF) ln Q
E = E° − (0.0591/n) log Q (at 25°C)
At equilibrium: E = 0, Q = K → E° = (0.0591/n) log K
As reactants are consumed and products build up, Q increases and E decreases. When the cell reaches equilibrium (Q = K), E = 0 — the cell is dead (no more current flows). For concentration cells (same electrodes, different concentrations), E° = 0 but E ≠ 0.
Molar conductivity (Λm) increases with dilution — more dissociation means more ions available to carry charge. For strong electrolytes, Kohlrausch's law: Λm = Λ°m − K√c (Λm increases linearly with √c as concentration decreases). For weak electrolytes, molar conductivity increases steeply at low concentration as dissociation nears completion. Λ°m for weak electrolytes (e.g., acetic acid) cannot be measured directly — calculated using Kohlrausch's law of independent migration of ions: Λ°m(CH₃COOH) = Λ°m(HCl) + Λ°m(CH₃COONa) − Λ°m(NaCl).
First Law: m = ZIt = (M/nF) × Q
Second Law: m₁/m₂ = E₁/E₂ (equivalent weights)
1 Faraday = 96500 C = charge of 1 mol electrons
First law: mass deposited ∝ charge passed (Q = It). Second law: masses deposited by same charge are in ratio of their equivalent weights. Applications: electroplating, electrorefining of copper, chlor-alkali process (NaCl → Cl₂ + H₂ + NaOH), anodising aluminium.
Corrosion is the electrochemical degradation of metals. Iron rusting is an electrochemical process: Fe acts as anode (oxidised: Fe → Fe²⁺ + 2e⁻) and dissolved oxygen acts as cathode (reduced: O₂ + 2H₂O + 4e⁻ → 4OH⁻). The Fe²⁺ and OH⁻ combine to form Fe(OH)₂, which oxidises further to rust (Fe₂O₃·xH₂O). Prevention methods: painting, galvanising (Zn coating — sacrificial anode), cathodic protection (connecting iron to Mg or Zn), alloying (stainless steel — Cr forms protective Cr₂O₃ layer).
📌 Daniel cell (primary): Zn-Cu, E° = 1.10 V
📌 Dry cell (Leclanché): Zn anode, MnO₂ cathode, NH₄Cl paste, ~1.5V
📌 Lead acid battery (secondary): Pb/PbSO₄|H₂SO₄|PbO₂/PbSO₄, 2V/cell, 6 cells=12V
📌 Nickel-cadmium: rechargeable, used in electronics
📌 Lithium-ion: high energy density, used in phones/EVs
📌 H₂-O₂ fuel cell: H₂ + ½O₂ → H₂O, only water as product, used in spacecraft
Electroplating deposits a thin layer of metal on another metal using electrolysis. The object to be plated is the cathode; the plating metal is the anode; electrolyte contains ions of the plating metal. Applications: silver plating cutlery, gold plating jewellery, chrome plating car parts, tin plating food cans (prevents corrosion), nickel plating for hardness. The thickness of deposit depends on current density and time (Faraday's first law).