Electrochemistry

Galvanic vs electrolytic

Galvanic cell: spontaneous redox → electrical energy. Anode oxidation (negative in many conventions for a cell supplying current), cathode reduction.

Electrolytic cell: external voltage drives non-spontaneous reaction. Anode is where oxidation still happens but it may be labelled positive in the electrolysis convention your book uses — follow your NCERT diagram labels.

Salt bridge or porous barrier keeps charge balance between half-cells. Cell notation: left anode | solution || solution | cathode right.

EMF and Nernst equation

Standard electrode potential E° is measured vs SHE. Cell E°_cell = E°_cathode − E°_anode for the written cell reaction.

Nernst equation at temperature T: E = E° − (RT/nF) ln Q. At 298 K, (RT/F) × 2.303 ≈ 0.0591 V so E = E° − (0.0591/n) log Q for log base 10 form in many Indian texts.

ΔG° = −nFE°_cell links thermodynamics and electrochemistry. K and E° connect at equilibrium when E = 0.

Conductance in solution

Conductivity κ (S cm^-1) depends on ion concentration and mobility. Molar conductivity Λ_m = κ × (1000/M) for mol L^-1. Strong electrolytes: Λ_m rises on dilution (inter-ionic forces drop). Weak electrolytes: ionisation increases on dilution so Λ_m rises sharply.

Kohlrausch law: Λ_m° for a salt = sum of limiting molar conductivities of ions. Used to find Λ_m° of weak acids from strong salt data.

Electrolysis applications

Faraday’s laws: mass deposited ∝ charge passed ∝ current × time. One faraday ≈ 96500 C mol^-1 of electrons.

Order of discharge at electrodes depends on electrode potential, concentration, and overpotential — syllabus usually gives simplified rules for aqueous vs molten electrolysis (e.g. Cu²⁺ vs H⁺ at cathode).