Exothermic reactions release energy, so why is the enthalpy change negative?

Because the sign describes what happens to the chemicals, not to you or to the surroundings. This single point of view fixes the whole topic.

Take the system's perspective

Everything in thermochemistry is measured from the point of view of the reacting substances — the system. The surroundings are everything else, including the beaker, the air and the thermometer.

Exothermic: the system loses energy to the surroundings. The system has less than it started with, so ΔH is negative. The surroundings gain that energy, so the thermometer reading goes up.

Endothermic: the system takes energy in from the surroundings. The system has more than it started with, so ΔH is positive. The surroundings lost it, so the temperature goes down.

The apparent contradiction — negative ΔH but a rising temperature — dissolves once you notice the two measurements are of different things.

Why it comes down to bonds

This is the explanation that earns marks rather than restating the definition.

Breaking bonds requires energy. Making bonds releases energy. Always, with no exceptions, and it is worth saying in those words.

Every reaction does both: old bonds break, new bonds form. Which way the overall energy goes depends only on which is bigger.

• Energy released making new bonds > energy needed breaking old ones → exothermic, ΔH negative
• Energy needed breaking old bonds > energy released making new ones → endothermic, ΔH positive

The common error is to think breaking bonds releases energy because an explosion involves things breaking apart. It does not. An explosion is exothermic because of the strong bonds in the products, not the weak ones in the reactants.

Reading the profile diagram

On an energy level diagram, look only at where the products sit relative to the reactants.

Products lower than reactants → energy was released → exothermic → ΔH is the downward arrow and is negative.

Products higher than reactants → energy was absorbed → endothermic → positive.

The hump in the middle is the activation energy, measured from the reactants up to the peak. It is not the enthalpy change and it is positive for both types, because every reaction needs some energy to get started. Mixing these two is the most common diagram mistake.

Worth memorising

Exothermic: combustion, neutralisation, most oxidation, respiration. Endothermic: thermal decomposition, photosynthesis, dissolving ammonium nitrate.

If a question describes a cold pack getting cold, it is telling you the reaction is endothermic and inviting you to say so with a positive ΔH.