3.1.5.5 - Catalysts

What a catalyst changes

A reaction only happens when particles collide with enough energy to get over the activation energy barrier. If too few particles can do that, the reaction is slow. A catalyst speeds the reaction up by changing that barrier, not by becoming part of the final products.

The key phrase is "alternative reaction route". A catalyst provides a different pathway from reactants to products, and that pathway has a lower activation energy than the uncatalysed route. Because the energy barrier is lower, a larger fraction of collisions can lead to reaction.

This does not mean the catalyst gives particles extra energy. The particles in the mixture still have the same range of energies at the same temperature. What changes is the minimum energy needed for success.

The diagram below compares uncatalysed and catalysed reaction profiles, making it clear that the catalyst lowers the activation energy without changing the overall energy difference between reactants and products.

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Using the Maxwell-Boltzmann distribution

For gas reactions, the Maxwell-Boltzmann distribution helps you picture why a catalyst works. The curve shows how many particles have each energy value at a particular temperature. Most particles have middling energies, while only a smaller fraction have very high energies.

The diagram below keeps the Maxwell-Boltzmann curve unchanged but moves the activation energy line to the left, so a larger area lies to the right of $E_a$.

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At constant temperature, adding a catalyst does not change the shape of the Maxwell-Boltzmann curve. The particles have not been heated, so their energy distribution is unchanged. Instead, the activation energy line moves to the left because the new route needs less energy.