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Snap a hand warmer and it heats up. Squeeze an instant cold pack and it turns icy. Both are chemical or physical processes that move energy, but in opposite directions. That direction is the difference between exothermic and endothermic.
The definitions
- An exothermic process releases heat to its surroundings. The surroundings get warmer. (Exo- means “outside”.)
- An endothermic process absorbs heat from its surroundings. The surroundings get colder. (Endo- means “inside”.)
The “system” is the reaction itself; the “surroundings” is everything else — the test tube, the water, your hand. When a beaker feels cold during a reaction, it’s because the reaction is pulling heat out of the beaker and your hand.
Enthalpy change, ΔH
Chemists measure the heat exchanged at constant pressure as the enthalpy change, ΔH, usually in kilojoules per mole (kJ/mol).
- Exothermic: ΔH is negative. The system loses energy.
- Endothermic: ΔH is positive. The system gains energy.
The sign is from the reaction’s point of view, which is why “releasing heat” gets a minus sign.
Burning methane: CH₄ + 2O₂ → CO₂ + 2H₂O ΔH ≈ −890 kJ/mol (exothermic)
Decomposing calcium carbonate (making lime): CaCO₃ → CaO + CO₂ ΔH ≈ +178 kJ/mol (endothermic — the kiln has to keep supplying heat)
Where the energy comes from: bonds
Every reaction breaks some bonds and makes new ones.
- Breaking a bond always absorbs energy. You have to pull the atoms apart.
- Making a bond always releases energy.
The overall ΔH is the balance between the two:
- If the new bonds are stronger than the old ones, more energy is released than absorbed → exothermic.
- If the new bonds are weaker, more energy goes in than comes out → endothermic.
A very common misconception is that “breaking bonds releases energy” — for example, that ATP releases energy by breaking a bond. It doesn’t; the energy comes from the stronger bonds formed afterwards.
Energy diagrams
Reaction energy diagrams plot energy as the reaction proceeds.
- In an exothermic diagram, the products sit lower than the reactants. The drop is ΔH.
- In an endothermic diagram, the products sit higher.
Both diagrams have a hump in between — the activation energy, the energy needed to get the reaction started. That’s why a match must be struck before it burns, even though burning is strongly exothermic. A catalyst lowers the hump without changing the start and end levels.
Everyday examples
Exothermic
- Combustion of any fuel: wood, gas, petrol, candle wax
- Respiration in your cells
- Neutralizing an acid with a base
- Iron rusting — the basis of disposable hand warmers, which use iron powder that oxidizes quickly
- Setting of concrete and plaster
- Condensation and freezing (physical changes: gas → liquid, liquid → solid)
- Dissolving calcium chloride or sodium hydroxide in water
Endothermic
- Photosynthesis (light supplies the energy)
- Thermal decomposition, like heating limestone
- Cooking an egg or baking bread
- Dissolving ammonium nitrate in water — the basis of instant cold packs
- Melting, evaporating and boiling (physical changes: solid → liquid → gas)
- Baking soda reacting with vinegar or citric acid (the mixture gets slightly cold)
Physical changes count too
Phase changes have enthalpy changes. Melting ice absorbs about 6.0 kJ/mol; evaporating water absorbs about 41 kJ/mol at 100 °C. That’s why sweating cools you: evaporating sweat takes heat from your skin. And it’s why steam burns are worse than boiling-water burns — steam releases that extra 41 kJ/mol as it condenses on your skin.
Endothermic but spontaneous?
If releasing energy is “downhill”, how can an endothermic process like ice melting on a warm day happen by itself? Because energy isn’t the only factor: nature also favours increasing entropy (roughly, disorder or spreading out). Ice melting and ammonium nitrate dissolving both increase entropy a lot, enough to make them go even while absorbing heat. The two factors are combined in a quantity called Gibbs free energy.
Quick answers
Is exothermic always spontaneous? Often, but not always — some exothermic reactions need a big push over their activation energy, and entropy matters too.
Why does ΔH have a negative sign for exothermic reactions? Because it’s measured from the system’s point of view: the system ends with less energy than it started with.
Is melting exothermic or endothermic? Endothermic. Freezing, its reverse, is exothermic.
Related
Temperature changes shift reversible reactions in predictable ways depending on whether they’re exothermic or endothermic — see Le Chatelier’s principle.
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