ElectrochemistryNEET Chemistry · Class 12 · NCERT Chapter 2

Electrochemistry typically gives 3 to 5 questions per NEET paper. The most common topics are standard electrode potential and cell EMF calculation (1 to 2 questions), the Nernst equation (1 question), and Faraday's laws (1 question). Conductance and batteries also appear occasionally. This makes it one of the highest-yield chapters in Class 12 Chemistry.

The standard electrode potential (E°) is the voltage of a half-cell measured under standard conditions (1 M concentration, 1 bar pressure, 25°C) relative to the Standard Hydrogen Electrode (SHE), which is assigned E° = 0.00 V. A positive E° means the half-cell is a stronger oxidising agent than hydrogen; it tends to accept electrons (get reduced). A negative E° means it is a stronger reducing agent than hydrogen; it tends to lose electrons (get oxidised). The more positive the E°, the stronger the oxidising agent.

In a galvanic cell (spontaneous), the anode is negative (oxidation occurs here; the electrode loses electrons) and the cathode is positive (reduction occurs here; the electrode gains electrons). In an electrolytic cell (non-spontaneous, driven by external power), the anode is connected to the positive terminal of the battery and the cathode to the negative terminal. The rule is: oxidation always occurs at the anode, and reduction always occurs at the cathode, in both types of cells.

The Nernst equation is E = E° - (0.0591/n) x log Q at 25°C, where E is the actual cell potential, E° is the standard cell potential, n is the number of electrons transferred, and Q is the reaction quotient (products/reactants using concentrations or pressures). You use it whenever concentrations differ from 1 M. A common NEET use is concentration cells, where both electrodes are the same metal but in solutions of different concentrations; E° = 0, so E = -(0.0591/n) x log Q. At equilibrium, E = 0 and Q = K, giving the relationship between E° and the equilibrium constant.

Conductance (G) is the ability of a conductor to allow current to flow; it is the reciprocal of resistance (G = 1/R) and is measured in siemens (S). Conductivity (kappa, k) is the conductance of a 1 cm cube of solution; it is measured in S cm-1 or S m-1. Molar conductivity (lambda_m) is the conductance of a solution containing 1 mole of electrolyte, with electrodes 1 cm apart, measured in S cm2 mol-1. As you dilute a solution, conductivity decreases (fewer ions per cm3) but molar conductivity increases (more volume per mole, so ions move more freely). At infinite dilution, molar conductivity reaches its maximum limiting value.

Faraday's first law states that the mass of substance deposited is directly proportional to the charge passed: m = Z x I x t, where Z is the electrochemical equivalent (Z = M / (n x F)), I is current in amperes, t is time in seconds, M is molar mass, n is the number of electrons per ion, and F = 96500 C/mol (Faraday's constant). Practical formula: m = (M x I x t) / (n x 96500). Faraday's second law states that for the same charge, the masses of different substances deposited are proportional to their equivalent masses (M/n). So if you pass the same charge through two different electrolytic cells, the ratio of masses deposited equals the ratio of their equivalent masses.

Kohlrausch's law states that at infinite dilution, the molar conductivity of an electrolyte equals the sum of the individual molar conductivities of its ions (lambda_m_infinity = v+ x lambda+_infinity + v- x lambda-_infinity, where v+ and v- are the number of cations and anions per formula unit). You use it to find the limiting molar conductivity of a weak electrolyte like acetic acid (CH3COOH), which cannot be measured directly (it would require infinite dilution in practice). You calculate it indirectly: lambda_m_infinity(CH3COOH) = lambda_m_infinity(HCl) + lambda_m_infinity(CH3COONa) - lambda_m_infinity(NaCl). Then you can find the degree of dissociation of the weak electrolyte: alpha = lambda_m / lambda_m_infinity.

A fuel cell is an electrochemical cell that converts chemical energy of a fuel (like hydrogen) into electrical energy continuously, as long as fuel and oxidant are supplied from outside. The reactants are not stored inside the cell. In a hydrogen-oxygen fuel cell, hydrogen is oxidised at the anode (H2 - 2e- gives 2H+) and oxygen is reduced at the cathode (O2 + 4H+ + 4e- gives 2H2O). The only product is water, making it environmentally clean. A battery, by contrast, contains a fixed amount of reactants sealed inside; once the reactants are used up, a primary battery is discarded or a secondary battery is recharged by reversing the reaction.