Energy changes in reactions
Every chemical reaction involves an energy change with its surroundings. In an exothermic reaction, energy is transferred to the surroundings, usually as heat, so the temperature rises. In an endothermic reaction, energy is taken in from the surroundings, so the temperature falls. Combustion, respiration and neutralisation are exothermic; thermal decomposition and photosynthesis are endothermic.
Key idea
The overall energy change is ΔH. Exothermic: ΔH is negative (energy released). Endothermic: ΔH is positive (energy absorbed).
Bond breaking and bond making
Breaking bonds in the reactants requires energy (endothermic step). Making new bonds in the products releases energy (exothermic step). If more energy is released making bonds than is used breaking them, the reaction is exothermic overall.
Worked example
H2 + Cl2 → 2HCl. Bond energies (kJ/mol): H–H = 436, Cl–Cl = 242, H–Cl = 431.
Energy in (bonds broken) = 436 + 242 = 678 kJ/mol.
Energy out (bonds made) = 2 × 431 = 862 kJ/mol.
ΔH = 678 − 862 = −184 kJ/mol. Negative, so the reaction is exothermic.
Energy level diagrams
On an energy level diagram, the reactants and products are drawn at different energies. For an exothermic reaction the products are lower than the reactants; for an endothermic reaction they are higher. The activation energy is the minimum energy colliding particles must have to react.
Remember
- Exothermic → heat out, temperature up, ΔH negative.
- Endothermic → heat in, temperature down, ΔH positive.
- Breaking bonds absorbs energy; making bonds releases energy.