Reaction Profiles and Activation Energy

In the last lesson you classified reactions as exothermic or endothermic from a temperature change. A reaction profile turns that idea into a diagram: it shows the energy of the chemicals as a reaction proceeds, from reactants on the left to products on the right. From a profile you can read off whether a reaction is exothermic or endothermic, see the energy "hill" the chemicals must climb to react — the activation energy — and understand how a catalyst speeds a reaction up. This lesson is part of Topic C3 of OCR Gateway Science A.

By the end of this lesson you should be able to draw and interpret reaction profiles for exothermic and endothermic reactions, define activation energy, explain how a catalyst works, and (Higher tier) calculate an overall energy change from bond energies.

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What a Reaction Profile Shows

A reaction profile (energy profile) is a graph with energy on the y-axis and progress of the reaction on the x-axis. It shows three things:

• the energy of the reactants (a level on the left);
• the energy of the products (a level on the right);
• the activation energy — the "hump" the chemicals must get over to react.

The overall energy change of the reaction is the difference in height between the reactant level and the product level.

Profile of an Exothermic Reaction

In an exothermic reaction the products are at a lower energy than the reactants, so energy is released to the surroundings. The product level sits below the reactant level.

Because the products are lower than the reactants, the overall energy change is negative (energy is given out). Combustion and neutralisation give profiles like this.

Profile of an Endothermic Reaction

In an endothermic reaction the products are at a higher energy than the reactants, because energy has been taken in from the surroundings. The product level sits above the reactant level.

Because the products are higher than the reactants, the overall energy change is positive (energy is taken in). Thermal decomposition gives a profile like this.

Exam Tip: Read the profile from left (reactants) to right (products). If the product level is lower than the reactant level it is exothermic; if it is higher it is endothermic. The "hump" is always there in both — that is the activation energy.

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Activation Energy

The activation energy ($E_a$Ea​) is the minimum energy that colliding particles must have for a reaction to happen. It is the height of the hump from the reactant level up to the top of the curve. Particles that collide with less than this energy simply bounce apart without reacting; only collisions with at least the activation energy are successful.

This is why some reactions need a spark or a flame to get started, even though they are exothermic overall: the spark gives the first few particles enough energy to get over the activation-energy hump, after which the energy released keeps the reaction going.

Activation energy and collision theory

Activation energy is best understood through collision theory. For two particles to react, they must (a) actually collide, and (b) collide with at least the activation energy — enough energy to start breaking the existing bonds. A collision with less than the activation energy is unsuccessful: the particles simply bounce apart unchanged. So the activation energy acts like a "toll" that every successful reaction must pay.

This explains why anything that gives particles more energy, or makes collisions more frequent, speeds a reaction up. Heating a reaction mixture, for example, gives the particles more kinetic energy, so a larger proportion of collisions have at least the activation energy and the reaction goes faster. The same logic explains why a reaction with a high activation energy is slow at room temperature — only a tiny fraction of collisions are energetic enough to react — and why lowering the activation energy (with a catalyst) makes such a dramatic difference to the rate.

Exam Tip: Define activation energy as the minimum energy needed for particles to react. Note that it appears on the profile of both exothermic and endothermic reactions — it is not the same as the overall energy change.

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How a Catalyst Works

A catalyst speeds up a reaction without being used up. It does this by providing a different reaction pathway that has a lower activation energy. On a reaction profile, a catalyst makes the hump lower — but it does not change the energy of the reactants or products, so the overall energy change stays exactly the same.

By lowering the activation energy, the catalyst means a larger proportion of collisions now have enough energy to be successful, so the reaction goes faster. The catalyst is not used up and is unchanged at the end.

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Higher Tier: Energy Change from Bond Energies