Exothermic and Endothermic Reactions Explained

The difference between exothermic and endothermic reactions is the direction energy flows. An exothermic reaction gives out energy to its surroundings, usually as heat, so the temperature rises: burning, neutralisation and respiration are all exothermic. An endothermic reaction takes energy in from its surroundings, so the temperature falls: an instant cold pack is the classic example. Which one a reaction is depends on its bonds: breaking bonds takes energy in, and making bonds gives energy out.

Why do reactions give out or take in energy?

Every chemical reaction breaks some bonds and makes new ones, and the energy change comes from the balance between the two.

• Breaking a bond is endothermic. The atoms attract each other, so it takes energy to pull them apart.
• Making a bond is exothermic. The same amount of energy is released when the atoms come together.

If the new bonds release more energy than it took to break the old ones, the extra is given out and the reaction is exothermic. If breaking the old bonds takes more than the new bonds give back, the reaction has to take the difference from its surroundings: endothermic.

The energy change of a reaction is called the enthalpy change, ΔH, measured in kJ/mol. The sign tells you the direction:

• ΔH negative: exothermic, energy given out, surroundings get hotter.
• ΔH positive: endothermic, energy taken in, surroundings get colder.

The sign looks backwards at first. It's from the point of view of the chemicals: in an exothermic reaction they lose energy, so ΔH is negative.

Examples of exothermic and endothermic reactions

Exothermic (ΔH negative)	Endothermic (ΔH positive)
Burning fuels: methane + oxygen	Thermal decomposition: calcium carbonate → calcium oxide + carbon dioxide
Neutralisation: acid + alkali	Photosynthesis: carbon dioxide + water → glucose + oxygen
Respiration in your cells	Dissolving ammonium nitrate in water (instant cold packs)
Iron rusting (hand warmers)	Citric acid + sodium hydrogencarbonate (sherbet fizzing)
Metals reacting with acids	Nitrogen + oxygen → nitrogen monoxide (in lightning)

Some figures: burning methane, the main gas in a cooker, gives out 890 kJ for every mole. Neutralising a strong acid with a strong alkali gives out about 57 kJ per mole of water formed. Breaking calcium carbonate down into lime takes in about 178 kJ/mol, which is why lime kilns run at over 800 °C.

Among the examples of endothermic reactions, photosynthesis is the most important. Plants take the energy from sunlight and store it in the bonds of glucose. Respiration releases it again, which makes respiration exothermic.

What is activation energy?

Activation energy (Ea) is the minimum energy that colliding particles need for a reaction to happen. Even an exothermic reaction has to break some bonds before it can make new ones, and that start-up cost has to be paid first.

That's why a gas cooker doesn't light itself. Methane and oxygen sit together quite happily until a spark gives a few molecules enough energy to react. Those reactions release heat, which gives the next molecules their activation energy, and the flame keeps itself going.

Only collisions with at least the activation energy succeed; the rest just bounce off. Raising the temperature makes particles move faster, so more collisions have enough energy. This is the idea behind collision theory, and it's why reactions speed up when they're heated.

How to read an energy profile diagram

An energy profile diagram (or reaction profile) shows the energy of the chemicals as the reaction goes from reactants to products.

• The reactants are a level on the left, and the products a level on the right.
• Between them is a hump. Its height above the reactants is the activation energy.
• The difference between the reactants and products is ΔH.

For an exothermic reaction the products are lower than the reactants: energy has been given out. For an endothermic reaction the products are higher: energy has been taken in. In both cases the hump comes first.

A catalyst gives the reaction a different route with a lower hump. More collisions now have enough energy, so the reaction is faster. The reactant and product levels don't move, so ΔH is exactly the same with or without a catalyst. Enzymes are biological catalysts that do this in your cells: see how temperature affects enzyme activity.

How to calculate enthalpy change from bond energies

To calculate enthalpy change from bond energies, add up the energy of every bond broken, add up the energy of every bond made, and subtract:

ΔH = total energy of bonds broken − total energy of bonds made

Worked example: hydrogen + chlorine

H₂ + Cl₂ → 2HCl, using average bond energies H–H = 436, Cl–Cl = 242 and H–Cl = 431 kJ/mol.

1. Bonds broken: one H–H and one Cl–Cl = 436 + 242 = 678 kJ
2. Bonds made: two H–Cl = 2 × 431 = 862 kJ
3. ΔH = 678 − 862 = −184 kJ/mol

Negative, so exothermic. A mixture of hydrogen and chlorine explodes in sunlight.

Worked example: nitrogen + oxygen

N₂ + O₂ → 2NO, using N≡N = 945, O=O = 498 and N=O = 630 kJ/mol.

1. Bonds broken: 945 + 498 = 1443 kJ
2. Bonds made: 2 × 630 = 1260 kJ
3. ΔH = 1443 − 1260 = +183 kJ/mol

Positive, so endothermic. Nitrogen's triple bond is one of the strongest there is, which is why the nitrogen and oxygen in air don't react unless something as hot as lightning or a car engine forces them to.

Bond-energy answers are estimates, because each value is an average over many different molecules. For methane burning they give about −818 kJ/mol, compared with the measured 890 kJ given out.

Common mistakes

• Subtracting the wrong way round. It's broken − made. Getting it backwards flips the sign.
• Forgetting to count every bond. CH₄ has four C–H bonds; 2H₂O has four O–H bonds.
• Saying a catalyst changes ΔH. It only lowers the activation energy.
• Thinking exothermic reactions need no energy to start. Most still need their activation energy first: a match, a spark or heat.
• Mixing up the sign. Exothermic is ΔH negative, even though the surroundings get hotter.

Frequently asked questions

Is burning exothermic or endothermic?

Burning (combustion) is always exothermic. The fuel reacts with oxygen to make carbon dioxide and water, and the very strong C=O and O–H bonds in the products release more energy than it took to break the bonds in the fuel and oxygen. It still needs a flame or spark to start, to supply the activation energy, but after that the heat it gives out keeps it going.

Is melting ice endothermic?

Yes. Melting takes in energy (334 J for every gram of ice) to break the forces holding the water molecules in their lattice, so it cools its surroundings, which is why ice cools a drink. It's a physical change rather than a chemical reaction, but the same rule applies: separating particles takes energy in. See the states of matter for more.

Why does a catalyst not change ΔH?

ΔH depends only on the energy of the reactants and the products, which a catalyst doesn't change: the same bonds are broken and the same bonds are made. A catalyst only provides a different route between them with a lower activation energy. It's like a tunnel through a hill instead of a road over the top: you start and finish at the same heights.

Try it yourself

Open Chemical Energetics and watch hydrogen and chlorine molecules collide at room temperature. Every collision is a dot on the energy profile climbing the activation-energy hill and rolling back. Press ⚡ Spark and watch the first molecules get over, then the heat they give out set off the rest. Then switch to N₂ + O₂, heat it to lightning temperatures and watch it cool itself down as it reacts. Turn on the catalyst to see the lower hump with ΔH unchanged. The same bonds are explored in the Molecule Builder.