Energy Changes in Chemical Reactions
Exothermic and Endothermic Reactions
All chemical reactions involve energy changes. Bonds in reactants are broken (requires energy) and bonds in products are formed (releases energy). The overall energy change determines whether the reaction is exothermic or endothermic.
Exothermic reaction:
- Energy released to the surroundings (as heat, light, sound)
- Temperature of surroundings increases
- Products have less energy than reactants
- Overall: energy released > energy absorbed (bond breaking vs bond forming)
Examples: Combustion (burning), oxidation (rusting), neutralisation, respiration, many displacement reactions.
Endothermic reaction:
- Energy absorbed from the surroundings
- Temperature of surroundings decreases
- Products have more energy than reactants
- Overall: energy absorbed > energy released
Examples: Thermal decomposition, dissolving ammonium nitrate in water, photosynthesis.
Reaction Profiles (Energy Diagrams)
A reaction profile (energy level diagram) shows the energy of reactants and products plotted against the progress of the reaction (reaction coordinate).
Exothermic reaction profile:
- Reactants are at a higher energy level than products
- ΔH (enthalpy change) is negative (energy is released)
- There is a hump above the reactant energy level — the activation energy (Eₐ)
Endothermic reaction profile:
- Products are at a higher energy level than reactants
- ΔH is positive (energy is absorbed)
- Activation energy is shown as the hump above the reactant level
Activation energy (Eₐ): The minimum energy required by colliding particles to start a reaction. A catalyst reduces the activation energy — more particles have enough energy to react → reaction is faster.
Bond Energy Calculations
During a chemical reaction:
- Bond breaking requires energy (endothermic step) — energy must be supplied to overcome the forces holding atoms together
- Bond forming releases energy (exothermic step) — energy is released as new bonds form
Overall energy change = Energy in (bonds broken) − Energy out (bonds formed)
- If energy in > energy out: endothermic overall (positive ΔH)
- If energy in < energy out: exothermic overall (negative ΔH)
Worked example — Hydrogen + Chlorine → Hydrogen Chloride: H₂ + Cl₂ → 2HCl
Bond energies (kJ/mol): H−H = 436; Cl−Cl = 242; H−Cl = 431
Energy in (bonds broken):
- 1 × H−H: 436 kJ
- 1 × Cl−Cl: 242 kJ
- Total in: 678 kJ
Energy out (bonds formed):
- 2 × H−Cl: 2 × 431 = 862 kJ
ΔH = 678 − 862 = −184 kJ/mol (exothermic — negative value means energy released)
Practical Applications
| Reaction type | Real-world use |
|---|---|
| Exothermic — combustion | Fuels (natural gas, petrol, coal) for heating and transport |
| Exothermic — neutralisation | Hand warmers (iron oxidation or calcium chloride + water) |
| Endothermic | Cold packs for sports injuries (ammonium nitrate + water) |
| Exothermic — respiration | All living organisms — releases energy for life processes |
WJEC Exam Tip
WJEC papers frequently combine bond energy data tables with questions asking you to:
- Calculate ΔH using bond energies
- Classify the reaction as exothermic or endothermic
- Explain why a catalyst changes reaction rate but not ΔH
Remember: A catalyst lowers activation energy but does not change the overall energy change (ΔH) of the reaction.
AI-generated · claude-opus-4-7 · v3-wjec-combined-science