Exothermic & Endothermic

Every chemical reaction involves an energy change. Heat energy is either given out to the surroundings or taken in from them — this is what splits reactions into two groups.

Exothermic reactions GIVE OUT heat energy to the surroundings.

• The temperature of the surroundings (the thermometer in the mixture, or your hand on the test tube) RISES.
• The enthalpy change ΔH is NEGATIVE (the chemicals lose energy to the surroundings).

Examples of exothermic reactions:

• Combustion — any fuel burning in oxygen (methane, ethanol, candle wax, petrol). This is why we use fuels for heat.
• Neutralisation — acid + alkali, e.g. HCl + NaOH → NaCl + H₂O. The mixture warms up by a few °C.
• Most displacement reactions — e.g. a more reactive metal pushing out a less reactive one: Zn + CuSO₄ → ZnSO₄ + Cu (mixture gets hot).
• Respiration — glucose + oxygen → carbon dioxide + water, releasing the energy that keeps living things warm and active.

Endothermic reactions TAKE IN heat energy from the surroundings.

• The temperature of the surroundings FALLS.
• The enthalpy change ΔH is POSITIVE (the chemicals gain energy from the surroundings).

Examples of endothermic reactions:

• Thermal decomposition — e.g. CaCO₃ → CaO + CO₂. You must keep heating it for the reaction to continue.
• Dissolving some salts — e.g. ammonium nitrate dissolving in water makes the water cold; this is how instant cold packs for sports injuries work.
• Photosynthesis — carbon dioxide + water → glucose + oxygen, using light energy taken in from the Sun.

Key idea — it is all about the SURROUNDINGS.

The words exothermic and endothermic describe what happens to the surroundings, which is what the thermometer measures:

Surroundings	Reaction type	ΔH	Mixture feels
Get HOTTER	EXOthermic	NEGATIVE	warm/hot
Get COLDER	ENDOthermic	POSITIVE	cold

An energy level diagram (also called a reaction profile) is a simple graph that shows how the energy of the chemicals changes during a reaction.

• X-axis = progress of reaction (left = reactants, right = products).
• Y-axis = energy.

Three things must be drawn/labelled correctly to earn full marks: the reactant and product energy levels, the activation energy (Ea) hump, and the enthalpy change (ΔH).

Exothermic profile.

• Reactants line is at a HIGHER energy.
• A peak (the "hump") rises above the reactants.
• Products line is at a LOWER energy than the reactants.
• ΔH arrow points DOWN (reactants → products) and ΔH is NEGATIVE.

Endothermic profile.

• Reactants line is at a LOWER energy.
• A peak rises above both reactants and products.
• Products line is at a HIGHER energy than the reactants.
• ΔH arrow points UP (reactants → products) and ΔH is POSITIVE.

How to label Ea and ΔH.

1. Activation energy (Ea) = the energy from the reactants line UP to the top of the peak. Draw the arrow from the reactants level to the peak.
2. Enthalpy change (ΔH) = the energy difference from reactants to products. Draw the arrow from the reactants level to the products level (down for exo, up for endo).

What activation energy means. Ea is the minimum energy that colliding particles must have before they can react. Even a very exothermic reaction needs this first "push" — for example, methane will not burn until a spark or flame supplies the activation energy. Once started, the heat given out keeps the reaction going.

Where does the energy change in a reaction actually come from? It comes from breaking the bonds in the reactants and making new bonds in the products.

Bond breaking is ENDOTHERMIC.

• A bond is a strong force of attraction holding atoms together.
• To break a bond you must pull the atoms apart, which needs energy IN from the surroundings.
• Breaking bonds therefore takes in energy = endothermic.

Bond making is EXOTHERMIC.

• When new bonds form, the atoms settle into a more stable, lower-energy arrangement.
• The leftover energy is given out to the surroundings.
• Making bonds therefore releases energy = exothermic.

Putting it together — the overall energy change.

In a reaction, old bonds break (energy in) AND new bonds make (energy out). The overall ΔH depends on which is bigger:

$\Delta H = (\text{energy to break bonds}) - (\text{energy released making bonds})$

• If more energy is released making new bonds than is taken in breaking old bonds → ΔH negative → exothermic.
• If more energy is taken in breaking bonds than is released making new bonds → ΔH positive → endothermic.

This is the deep reason exothermic and endothermic reactions exist — it all comes down to the balance between bonds broken and bonds made.

In the lab you can tell which type a reaction is just by watching the thermometer:

• Temperature goes UP → heat given out → exothermic.
• Temperature goes DOWN → heat taken in → endothermic.

Typical exam scenario. A student mixes two solutions in a polystyrene cup and records the temperature before and after.

• If the temperature rises (e.g. 21 °C → 28 °C), the reaction is exothermic, ΔH is negative.
• If the temperature falls (e.g. 21 °C → 14 °C), the reaction is endothermic, ΔH is positive.

Why a polystyrene cup? Polystyrene is a poor conductor (a good insulator), so it reduces heat loss/gain to the surroundings — giving a more accurate temperature change.

Linking it all back to bonds. A reaction that releases heat (exothermic) does so because making the new bonds in the products gives out more energy than was needed to break the old bonds in the reactants. A reaction that takes heat in (endothermic) does the opposite.

Quick checklist for any reaction question:

1. Does the temperature rise or fall? → exo or endo.
2. What is the sign of ΔH? → negative (exo) or positive (endo).
3. On an energy diagram, are the products lower or higher than the reactants? → lower (exo) or higher (endo).
4. Which bonds dominate? → bonds made release more (exo) or bonds broken take more (endo).