Exothermic Reactions
An exothermic reaction releases energy into the surroundings — the surroundings (including the test tube or beaker) get warmer.
More energy is released when new bonds form than is spent breaking old bonds.
Examples:
- Combustion (burning fuel) — releases heat and light
- Neutralisation of an acid with an alkali — test tube warms up
- Respiration — cells release energy from glucose
- Explosions
- Hand warmers (iron oxidising slowly)
Endothermic Reactions
An endothermic reaction absorbs energy from the surroundings — the surroundings get cooler.
More energy is required to break bonds than is released when new bonds form.
Examples:
- Thermal decomposition of carbonates — requires continuous heating
- Photosynthesis — absorbs light energy
- Dissolving ammonium nitrate in water — the solution feels cold
- Some sports injury cold packs work this way
A quick test: place a thermometer in the reaction mixture. If the temperature rises → exothermic. If it falls → endothermic.
Activation Energy
All reactions need a small amount of energy to start — this is the activation energy. Think of it as the energy needed to overcome an initial barrier. Matches need to be struck; fuels need a flame or spark.
Catalysts lower the activation energy needed, making reactions happen more easily and at lower temperatures.
Energy Profile Diagrams
An energy profile diagram shows how energy changes from reactants to products:
- Exothermic: products are at a lower energy level than reactants — energy goes into the surroundings
- Endothermic: products are at a higher energy level — energy is taken from the surroundings
Memory tip: "EXothermic = EXits energy (heat released)". "ENdothermic = ENters the reaction (heat absorbed)".
Real world: Hand warmers (exothermic)
Disposable hand warmers contain iron powder, salt, and water in a permeable pouch. When the pack is opened and shaken, oxygen from the air reacts with the iron in an exothermic slow oxidation reaction: 4Fe + 3O₂ → 2Fe₂O₃. The energy released heats the pack to around 50–60 °C for several hours. The same chemistry that causes iron to rust is harnessed to keep your hands warm in winter.
Real world: Sports cold packs (endothermic)
Instant cold packs used in sports first aid contain a pouch of water surrounded by ammonium nitrate crystals. When you crack the inner pouch, the ammonium nitrate dissolves in the water. This dissolving is endothermic — it absorbs heat from the surroundings, cooling the pack to around 2–4 °C within seconds. The same principle of endothermic energy absorption creates the cooling effect.
Real world: Cooking with exothermic combustion
Every gas cooker uses the exothermic combustion of methane: CH₄ + 2O₂ → CO₂ + 2H₂O + energy. The energy released heats your pan and cooks your food. Conversely, cooking itself requires endothermic processes — dissolving sugar into a hot syrup, for example, absorbs heat. Understanding exothermic and endothermic reactions explains why cooking works the way it does.
Where do you see this in real life? Hand warmers and fires (exothermic — releases heat), cold packs and sweating (endothermic — absorbs heat), car engines burning petrol (exothermic), photosynthesis in plants (endothermic) — energy changes in chemical reactions power everything from your body to your kitchen to your car.