Energy changes in chemical reactions

Exothermic and endothermic reactions

Exothermic reaction: releases energy to surroundings → temperature INCREASES.

• Products have LESS energy than reactants.
• Examples: combustion, neutralisation, respiration, hand warmers, many oxidation reactions.

Endothermic reaction: takes in energy from surroundings → temperature DECREASES.

• Products have MORE energy than reactants.
• Examples: photosynthesis, thermal decomposition (CaCO₃ → CaO + CO₂), dissolving ammonium nitrate, sports cold packs.

Reaction profile diagrams (energy level diagrams)

For an exothermic reaction:

• Reactants are at a HIGHER energy level than products.
• Curve rises to a peak (the activation energy, Eₐ) then falls to product level.
• Energy difference = enthalpy change (ΔH), which is NEGATIVE for exothermic.

For an endothermic reaction:

• Reactants at a LOWER energy level than products.
• Curve rises and stays at a higher level.
• ΔH is POSITIVE for endothermic.

The activation energy (Eₐ) is the minimum energy particles must have to react — height from reactant energy level to the peak.

A catalyst lowers Eₐ (shown as a lower peak) → more particles have enough energy → faster rate.

Bond energies

During a chemical reaction:

• Bond breaking: requires energy input (endothermic).
• Bond making: releases energy (exothermic).

Overall energy change = energy to break bonds in reactants − energy released making bonds in products

If energy IN (breaking) > energy OUT (making) → endothermic. If energy IN (breaking) < energy OUT (making) → exothermic.

Bond energy calculation example: H₂ + Cl₂ → 2HCl Bond energies: H−H = 436 kJ/mol; Cl−Cl = 243 kJ/mol; H−Cl = 432 kJ/mol

Energy in (breaking): 436 + 243 = 679 kJ Energy out (making): 2 × 432 = 864 kJ Overall: 679 − 864 = −185 kJ (negative = exothermic)