Energetics (C3.2)

Energy changes in chemical reactions appear on almost every Gateway A Chemistry paper. Bond energy calculations and reaction profile diagrams are the most common exam items.

Exothermic and endothermic reactions

All chemical reactions involve breaking bonds (requires energy) and making new bonds (releases energy). The overall energy change depends on which dominates.

Exothermic reaction: energy released to surroundings > energy absorbed.

• Temperature of surroundings increases.
• Products have LESS energy than reactants.
• ΔH is negative.
• Examples: combustion, neutralisation, oxidation, respiration, hand warmers.

Endothermic reaction: energy absorbed from surroundings > energy released.

• Temperature of surroundings decreases.
• Products have MORE energy than reactants.
• ΔH is positive.
• Examples: thermal decomposition, photosynthesis, dissolving ammonium nitrate in water, sports cold packs.

Reaction profile (energy diagram)

A reaction profile plots energy (y-axis) against progress of reaction (x-axis):

• Exothermic: reactants sit higher than products. The gap = energy released = ΔH (negative).
• Endothermic: reactants sit lower than products. The gap = energy absorbed = ΔH (positive).
• The hump between reactants and products = activation energy (Eₐ) — the minimum energy needed to start the reaction.
• A catalyst lowers the activation energy hump but does not change the reactant or product energy levels, and does not change ΔH.

Bond energies

To calculate the overall energy change:

ΔH = Energy needed to break bonds (endothermic) − Energy released forming bonds (exothermic)

• If the result is negative → exothermic (more energy released than absorbed).
• If the result is positive → endothermic (more energy absorbed than released).

Bond energy values (supplied in exam):

Common values at GCSE:

• H–H: 436 kJ/mol
• O=O: 498 kJ/mol
• O–H: 464 kJ/mol
• H–Cl: 431 kJ/mol
• Cl–Cl: 243 kJ/mol
• C–H: 413 kJ/mol
• C=O: 805 kJ/mol (in CO₂)

Worked example: formation of water from H₂ + O₂

2H₂ + O₂ → 2H₂O

Bonds broken:

• 2 × H–H: 2 × 436 = 872 kJ
• 1 × O=O: 498 kJ
• Total in = 1,370 kJ

Bonds formed:

• 4 × O–H: 4 × 464 = 1,856 kJ
• Total out = 1,856 kJ

ΔH = 1,370 − 1,856 = −486 kJ/mol → exothermic ✓

Practical measurement of energy change

Calorimetry — simple version:

1. Dissolve/react a measured amount of substance in a known volume of water.
2. Record temperature change ΔT.
3. Q = m × c × ΔT (where c for water = 4.18 J/g/°C, m in grams).
4. Divide by moles to get ΔH per mole.

Errors: heat loss to surroundings, incomplete combustion, evaporation.

Common Gateway-paper mistakes

1. Confusing exothermic (energy out, temperature rises) with endothermic.
2. Drawing a reaction profile with products higher than reactants for an exothermic reaction (wrong).
3. Forgetting to multiply bond energies by the number of bonds in the equation.
4. Saying a catalyst changes the energy of reactants or products — it only lowers the activation energy hump.
5. Getting the ΔH formula the wrong way round: it's bonds broken MINUS bonds formed (not the other way).