Energy changes in chemical reactions

Exothermic and endothermic reactions

Every chemical reaction involves an energy change, usually as heat.

Exothermic reactions release energy to the surroundings — the temperature of the surroundings (e.g. the solution or air around the reaction) increases.

• Examples: combustion, neutralisation, respiration, many oxidation reactions, adding metals to acids.
• ΔH is negative (energy released): ΔH < 0.

Endothermic reactions absorb energy from the surroundings — the temperature of the surroundings decreases (feels cold).

• Examples: thermal decomposition, photosynthesis, dissolving ammonium nitrate in water, the reaction of citric acid with sodium hydrogen carbonate.
• ΔH is positive (energy absorbed): ΔH > 0.

Energy level diagrams (reaction profiles):

• Exothermic: products are at a lower energy level than reactants. The "hump" = activation energy (Eₐ). ΔH = energy of products − energy of reactants (negative).
• Endothermic: products are at a higher energy level than reactants. ΔH is positive.

Bond energies (bond enthalpies)

Chemical reactions involve breaking bonds in reactants (requires energy input) and making bonds in products (releases energy). The overall energy change is the difference.

Bond energy (bond enthalpy): the energy required to break one mole of a particular bond in gaseous molecules (kJ mol⁻¹). It is always positive (breaking bonds is always endothermic).

Calculating ΔH from bond energies: ΔH = Σ(bonds broken) − Σ(bonds made)

• If positive: more energy used breaking bonds than released making → endothermic overall.
• If negative: more energy released making bonds than used breaking → exothermic overall.

Example: H₂ + Cl₂ → 2HCl Bond energies: H−H = 436, Cl−Cl = 242, H−Cl = 431 kJ mol⁻¹ Bonds broken: 1 × H−H (436) + 1 × Cl−Cl (242) = 678 kJ Bonds made: 2 × H−Cl (2 × 431) = 862 kJ ΔH = 678 − 862 = −184 kJ mol⁻¹ (exothermic)

Note: bond energy calculations give approximate ΔH values because they use average bond energies — actual values depend on the molecule's environment.

Calorimetry — measuring energy changes

Simple calorimetry (WJEC required practical):

• Measure a known volume of solution or mass of water in a polystyrene cup (good insulator).
• Record initial temperature.
• Add the reactant (e.g. dissolve a salt, mix acid and base, or burn a fuel under the cup).
• Record the maximum (or minimum for endothermic) temperature.
• Calculate:

q = mcΔT

Where:

• q = energy change (J or kJ)
• m = mass of solution/water (g) — assume density = 1 g/cm³ for aqueous solutions
• c = specific heat capacity of water = 4.18 J g⁻¹ °C⁻¹
• ΔT = temperature change (°C)

To find molar enthalpy change (ΔH in kJ mol⁻¹): ΔH = −q ÷ n (where n = moles of limiting reagent; negative sign for exothermic)

Sources of error in calorimetry:

• Heat lost to surroundings / poor insulation
• Heat absorbed by the calorimeter itself
• Evaporation of solvent
• Incomplete combustion (for spirit lamp calorimetry)

Fuel energy comparison

Spirit lamp calorimetry compares fuels by burning a known mass under a copper calorimeter of water. Energy per gram = q ÷ mass of fuel burned. This underestimates the true value due to heat losses.

Common examiner traps

1. ΔH sign convention: exothermic is negative (energy released from system to surroundings); endothermic is positive. Students often get these reversed.
2. Bond breaking is endothermic; bond making is exothermic — the net effect determines overall ΔH.
3. ΔH = bonds broken − bonds made (not made − broken). A positive result means more energy was put in → endothermic.
4. q = mcΔT uses mass of solution — not mass of the solid added (which is usually negligible or the question will specify).
5. Activation energy is not ΔH: Eₐ is the energy barrier to start the reaction; ΔH is the net energy change overall.