CC4.2 — Energy changes (Edexcel 1SC0)

Exothermic and endothermic reactions

Exothermic: energy released to surroundings → temperature of surroundings increases. Examples: combustion, neutralisation, respiration, oxidation of metals.

Endothermic: energy absorbed from surroundings → temperature of surroundings decreases. Examples: thermal decomposition, photosynthesis, dissolving ammonium nitrate.

Reaction profiles (energy diagrams)

• Shows energy of reactants and products.
• Activation energy (Ea): minimum energy needed for a reaction to occur (the "energy hill").
• Exothermic: products have lower energy than reactants (ΔH < 0).
• Endothermic: products have higher energy than reactants (ΔH > 0).

Bond energy calculations

Breaking bonds: endothermic (energy in). Forming bonds: exothermic (energy out).

$$\Delta H = \sum \text{bond energies broken} - \sum \text{bond energies formed}$$

If ΔH < 0: exothermic. If ΔH > 0: endothermic.

Example: H₂ + Cl₂ → 2HCl

• Bonds broken: H-H (436 kJ/mol) + Cl-Cl (243 kJ/mol) = 679 kJ
• Bonds formed: 2 × H-Cl (432 kJ/mol) = 864 kJ
• ΔH = 679 − 864 = −185 kJ/mol (exothermic)