Energy Changes in Chemical Reactions

Exothermic reactions

In an exothermic reaction, energy is released to the surroundings, so the temperature of the surroundings increases. The products have less energy than the reactants.

Common examples:

• Combustion of fuels (burning methane, wood, coal).
• Neutralisation of acids with alkalis.
• Respiration: C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + energy.
• Oxidation of metals (rusting — slow exothermic).
• Hand warmers (iron oxidation or chemical crystallisation).

Energy change symbol: ΔH < 0 (negative — energy is lost from the reaction system).

Endothermic reactions

In an endothermic reaction, energy is absorbed from the surroundings, so the temperature of the surroundings decreases. The products have more energy than the reactants.

Common examples:

• Thermal decomposition (e.g. CaCO₃ → CaO + CO₂; requires a lot of heat input).
• Photosynthesis: 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂ (absorbs light energy).
• Dissolving ammonium nitrate in water (cold packs).
• Electrolysis.

Energy change symbol: ΔH > 0 (positive — energy is absorbed into the reaction system).

Energy profile diagrams

An energy profile (reaction coordinate diagram) shows the energy of particles as the reaction proceeds.

Exothermic profile:

• Reactants higher energy than products.
• ΔH = energy of products − energy of reactants = negative value.
• Activation energy (Eₐ) = height of the energy "hump" above reactants.

Endothermic profile:

• Products higher energy than reactants.
• ΔH = positive value.
• Eₐ = height of hump above reactants (always positive — energy always needed to start a reaction).

Effect of a catalyst on the energy profile:

• Lowers the activation energy (the hump is lower).
• ΔH is unchanged (the overall energy change is the same — start and end energies are the same).
• There may be two smaller humps if the mechanism has two steps.

Bond breaking and bond making

All chemical reactions involve:

• Breaking bonds in reactants — this requires energy input (endothermic step).
• Making bonds in products — this releases energy (exothermic step).

Overall energy change: ΔH = energy needed to break bonds − energy released when bonds form

If more energy is released in bond forming than absorbed in bond breaking → overall exothermic (ΔH < 0). If more energy is absorbed in bond breaking → overall endothermic (ΔH > 0).

Bond energy calculation: Use tabulated bond energies (kJ/mol) to calculate ΔH.

Example: H₂ + Cl₂ → 2HCl Bond energies: H–H = 436, Cl–Cl = 242, H–Cl = 431 kJ/mol. Energy to break: 436 + 242 = 678 kJ. Energy released: 2 × 431 = 862 kJ. ΔH = 678 − 862 = −184 kJ/mol (exothermic).

CCEA practical — measuring enthalpy of neutralisation

Measure temperature change when acid and alkali are mixed in a polystyrene cup (good insulator). Use q = mcΔT to calculate energy released (m = mass of solution, c = 4.18 J/g/°C for dilute aqueous, ΔT = temperature rise).