Everything on the GCSE Chemistry Energy Changes poster is written out below, section by section. Use it to search the sheet, copy parts into your own notes, or check a fact quickly.
Exothermic vs endothermic
All reactions are either exothermic (release energy) or endothermic (absorb energy).
Exothermic reactions release energy to the surroundings, so the temperature of the surroundings rises. Examples include combustion, neutralisation, oxidation, and hand warmers.
Endothermic reactions take in energy from the surroundings, so the temperature of the surroundings falls. Examples include thermal decomposition, citric acid + sodium hydrogencarbonate, and sports cold packs.
Reaction profiles
A reaction profile plots energy against the progress of the reaction. The initial rise in the profile is the activation energy (Ea) - the minimum energy reactants need to collide and react.
- Exothermic profile: products are at a lower energy level than reactants. The overall energy change is negative.
- Endothermic profile: products are at a higher energy level than reactants. The overall energy change is positive.
In both cases the curve rises to a peak (the transition state) before falling to the product energy level. The activation energy is measured from the reactants' energy level up to that peak.
Bond energies
- Breaking bonds is endothermic - energy is supplied.
- Making bonds is exothermic - energy is released.
Overall energy change = energy supplied to break reactant bonds - energy released to form product bonds
- A positive value (+) means more energy was taken in than released, so the reaction is endothermic.
- A negative value (-) means more energy was released than taken in, so the reaction is exothermic.
Worked example: H₂ + Cl₂ → 2HCl
| Bond | Bond energy (kJ/mol) |
|---|---|
| H-H | 436 |
| Cl-Cl | 243 |
| H-Cl | 431 |
Breaking bonds: H-H = 436, Cl-Cl = 243. Total = 436 + 243 = 679 kJ/mol
Forming bonds: 2 × H-Cl = 2 × 431 = 862 kJ/mol
Overall energy change = 679 - 862 = -183 kJ/mol - this is exothermic.