Energy changes — section overview
Section C5 covers exothermic and endothermic reactions, bond energies and the energy changes in chemical reactions.
Exothermic vs endothermic reactions
| Type | Energy transfer | Temperature change | Products vs reactants |
|---|---|---|---|
| Exothermic | Heat released to surroundings | Temperature rises | Lower energy than reactants |
| Endothermic | Heat absorbed from surroundings | Temperature falls | Higher energy than reactants |
Exothermic examples: combustion, neutralisation, respiration, hand warmers. Endothermic examples: thermal decomposition, photosynthesis, sports cold packs.
Energy profile diagrams
In an energy profile diagram:
- Activation energy (Ea): energy needed to start a reaction (the "hill"); the higher the hill, the slower the reaction
- Exothermic: products lower than reactants; ΔH is negative
- Endothermic: products higher than reactants; ΔH is positive
Bond energies (HT)
Chemical reactions involve breaking bonds (requires energy — endothermic) and forming bonds (releases energy — exothermic).
$$\Delta H = \text{Bond energies broken} - \text{Bond energies formed}$$
- Negative ΔH → more energy released (forming) than absorbed (breaking) → exothermic
- Positive ΔH → more energy absorbed (breaking) than released (forming) → endothermic
Reaction profiles and catalysts
A catalyst lowers the activation energy without being used up. On the energy profile, the "hill" is lower — same start and end energies.
Cells and batteries (HT)
Chemical reactions in cells transfer energy as electrical energy. The further apart the two metals in the reactivity series, the greater the voltage produced.
Common exam mistakes in C5
- Exothermic reactions release energy — temperature of surroundings increases (the thermometer shows higher temperature)
- ΔH for exothermic is NEGATIVE — examiners expect −kJ/mol not +kJ/mol
- Bond breaking absorbs energy; bond forming releases energy — students often reverse this
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