Energetics

Conservation of energy

• Energy cannot be created or destroyed during a chemical reaction — only transferred between the reaction mixture (the system) and its surroundings
• When energy leaves the system, the surroundings warm up by the same amount
• When energy enters the system, the surroundings cool down by the same amount
• A thermometer placed in the reaction mixture is the simplest way to detect this transfer

Exothermic reactions

• An exothermic reaction transfers energy out of the system to the surroundings
• The temperature of the reaction mixture, together with any solvent it sits in, rises
• The chemical energy stored in the products is lower than in the reactants — the difference is what escapes as heat
• The enthalpy change is negative: ΔH < 0
• Reactions that are always exothermic include:
  • Combustion of fuels
  • Neutralisation of an acid by a base
  • The reaction of a reactive metal with water or acid
  • Many displacement and oxidation reactions

Endothermic reactions

• An endothermic reaction transfers energy into the system from the surroundings
• The temperature of the reaction mixture falls
• The products store more chemical energy than the reactants
• The enthalpy change is positive: ΔH > 0
• Reactions that are commonly endothermic include:
  • Thermal decomposition of metal carbonates
  • The reaction between citric acid and sodium hydrogencarbonate
  • The dissolving of some ionic salts (e.g. ammonium nitrate) — useful in cold packs
  • Photosynthesis (energy supplied by sunlight)

Reading the experimental data

• A rise in temperature = exothermic; a fall = endothermic. Mark schemes credit the direction of the temperature change plus the named conclusion