Thermochemistry & Hess's Law: JEE Main Complete Guide

Thermochemistry is the quantitative backbone of thermodynamics in Chemistry, contributing two to four marks per JEE Main session. The chapter tests enthalpy calculations using Hess's law, bond enthalpies, and Kirchhoff's equation. Once you master the sign conventions and the algebraic manipulation of thermochemical equations, this chapter becomes one of the most reliable scoring zones in the entire Chemistry paper.

Enthalpy and Sign Conventions

Enthalpy change ΔH = H_products − H_reactants. For exothermic reactions (heat released), ΔH < 0. For endothermic reactions (heat absorbed), ΔH > 0. Standard enthalpy ΔH° is measured at 25°C (298 K) and 1 atm. Standard enthalpy of formation ΔH°f is the enthalpy change when one mole of a compound is formed from its elements in their standard states — by definition, ΔH°f of any element in its standard state is zero. The standard enthalpy of a reaction is ΔH°rxn = Σ ΔH°f(products) − Σ ΔH°f(reactants). This relation is the most direct calculation tool for reaction enthalpies.

Enthalpy of combustion is always negative (exothermic). Enthalpy of neutralisation of a strong acid with a strong base is always approximately −57.1 kJ/mol (enthalpy of formation of water from H⁺ and OH⁻). Enthalpy of neutralisation of a weak acid is less negative because energy is consumed to ionise the weak acid. JEE uses this difference to test understanding of why enthalpy of neutralisation varies. Take a free thermochemistry mock to check your calculation accuracy.

Hess's Law and Its Applications

Hess's law: since enthalpy is a state function, ΔH for a reaction is independent of the path taken. Practically, this means you can add and subtract thermochemical equations algebraically to obtain the enthalpy of a target reaction. When adding equations, multiply the ΔH values by the same scalar used to balance the equation. When reversing an equation, reverse the sign of ΔH. JEE typically gives three or four equations and asks for the enthalpy of a target reaction — the skill is identifying which equations to add, which to reverse, and what multipliers to use. Setting up the algebra carefully before calculating is faster than trial and error.

Bond enthalpy method: ΔH_rxn ≈ Σ (bond enthalpies of bonds broken) − Σ (bond enthalpies of bonds formed). Breaking bonds requires energy (endothermic); forming bonds releases energy (exothermic). This method gives approximate values for gas-phase reactions. JEE gives a bond enthalpy table and asks for the enthalpy of a specific reaction, requiring you to identify which bonds break and which form. Draw the Lewis structures of reactants and products to count bonds correctly — the most common error is missing a bond in a structure.

Kirchhoff's Equation and Entropy

Kirchhoff's equation adjusts enthalpy for temperature changes: ΔH(T₂) = ΔH(T₁) + ΔCₚ(T₂ − T₁), where ΔCₚ = Σ Cₚ(products) − Σ Cₚ(reactants). JEE occasionally tests this when asking how enthalpy changes between 25°C and another temperature. Entropy change ΔS = Σ S°(products) − Σ S°(reactants) in J/mol·K. Gibbs free energy ΔG = ΔH − TΔS determines spontaneity: negative ΔG means spontaneous. The four ΔH–ΔS combination cases for spontaneity are directly tested (ΔH < 0, ΔS > 0 → always spontaneous; ΔH > 0, ΔS < 0 → never spontaneous; mixed signs → temperature-dependent). Our equilibrium guide connects ΔG to equilibrium constants via ΔG° = −RT ln K, tying thermochemistry to chemical equilibrium in a complete physical chemistry framework. See also our chemistry strategy guide for how to schedule thermochemistry in your overall plan.