What you'll learn
Energy changes accompany every chemical reaction: some give out heat, others take it in. In this guide you will learn the difference between exothermic and endothermic reactions, how to interpret reaction profiles and activation energy, the meaning of bond breaking and bond making in terms of energy, how to calculate energy changes from bond energies, and some everyday applications such as hand warmers and sports cooling packs. These ideas underpin rates of reaction and industrial chemistry.
Key terms and definitions
Exothermic reaction — a reaction that transfers energy to the surroundings (temperature rises).
Endothermic reaction — a reaction that takes in energy from the surroundings (temperature falls).
Activation energy — the minimum energy needed for a reaction to start.
Reaction profile — a diagram showing the energy of reactants and products through a reaction.
Bond energy — the energy needed to break (or released when making) one mole of a particular bond.
Conservation of energy — energy is not created or destroyed, only transferred.
Core concepts
Exothermic and endothermic reactions
In an exothermic reaction, energy is transferred to the surroundings, usually as heat, so the temperature of the surroundings rises. Examples include combustion, neutralisation, and many oxidation reactions. In an endothermic reaction, energy is taken in from the surroundings, so the temperature falls. Examples include thermal decomposition and the reaction of citric acid with sodium hydrogen carbonate.
Reaction profiles
A reaction profile shows how energy changes during a reaction. For an exothermic reaction, the products are at a lower energy than the reactants, so energy is released; the overall energy change is negative. For an endothermic reaction, the products are at a higher energy than the reactants. In both cases there is an energy "hump" between reactants and products — the activation energy, the minimum energy needed for the reaction to start (to break the initial bonds).
Activation energy
The activation energy is the minimum energy that colliding particles must have to react. A higher activation energy means fewer collisions are successful, so the reaction is slower. A catalyst works by providing an alternative pathway with a lower activation energy, speeding up the reaction.
Bond breaking and bond making
Chemical reactions involve breaking bonds in the reactants and making new bonds in the products:
- Breaking bonds is endothermic — energy must be supplied.
- Making bonds is exothermic — energy is released.
If more energy is released making bonds than is used breaking them, the reaction is exothermic. If more energy is needed to break bonds than is released making them, the reaction is endothermic.
Calculating energy change from bond energies
The overall energy change can be calculated as:
Energy change = (energy to break bonds in reactants) − (energy released making bonds in products)
A negative value means exothermic; a positive value means endothermic. Use the bond energies given in the question.
Everyday applications
Exothermic reactions are used in self-heating cans and hand warmers. Endothermic reactions are used in instant cold packs for sports injuries, which take in heat and feel cold.
Worked examples
Example 1: Classifying a reaction
A reaction causes the temperature of the surroundings to fall. Is it exothermic or endothermic?
A fall in temperature means heat is taken in from the surroundings, so it is endothermic.
Example 2: Bond energy calculation
For a reaction, breaking bonds needs 1000 kJ and making bonds releases 1200 kJ. What is the energy change?
Energy change = 1000 − 1200 = −200 kJ. The negative value means the reaction is exothermic.
Example 3: Role of a catalyst
How does a catalyst affect the reaction profile?
A catalyst lowers the activation energy (a smaller "hump") by providing an alternative pathway, so the reaction is faster, but the overall energy change is unchanged.
Common mistakes and how to avoid them
Mixing up exothermic and endothermic. Exothermic = energy out, surroundings warm; endothermic = energy in, surroundings cool.
Getting bond energy logic backwards. Breaking bonds takes in energy; making bonds releases energy.
Sign errors in calculations. Energy change = bonds broken − bonds made; negative = exothermic.
Forgetting activation energy. Both reaction types have an activation energy "hump" at the start.
Saying a catalyst changes the energy change. It only lowers activation energy, not the overall energy change.
Exam technique for Energy Changes
State the direction of energy transfer and the temperature change for the surroundings.
Draw and label reaction profiles, marking reactants, products, activation energy and the overall energy change.
Apply the bond energy rule — breaking is endothermic, making is exothermic.
Calculate energy changes as bonds broken minus bonds made, with the correct sign.
Link catalysts to lower activation energy, not to the overall energy change.
Quick revision summary
Every reaction has an energy change. Exothermic reactions release energy to the surroundings (temperature rises) — combustion, neutralisation, oxidation. Endothermic reactions take in energy (temperature falls) — thermal decomposition. On a reaction profile, exothermic reactions show products lower than reactants (negative overall change) and endothermic reactions show products higher; both have an activation energy "hump" — the minimum energy to start the reaction. At the molecular level, breaking bonds is endothermic (energy supplied) and making bonds is exothermic (energy released); the reaction is exothermic if more energy is released making bonds than is used breaking them. Calculate the overall change as bonds broken − bonds made (negative = exothermic). A catalyst lowers the activation energy by offering an alternative pathway, speeding up the reaction without changing the overall energy change. Everyday uses include exothermic hand warmers and endothermic cold packs. Classify reactions correctly, draw and label profiles, apply the bond-energy rule, and watch your signs in calculations.