How does a catalyst affect the rate of a reaction?

A catalyst increases the rate of a reaction by providing an alternative reaction pathway with a lower activation energy.

In more detail, a catalyst is a substance that can speed up a chemical reaction without being consumed or altered in the process. It achieves this by lowering the activation energy, which is the minimum energy required for a reaction to occur. The catalyst provides an alternative reaction pathway that requires less energy, allowing the reaction to proceed more quickly.

The way a catalyst works can be visualised using an energy profile diagram. In a reaction without a catalyst, the reactants must overcome a high energy barrier, represented by the activation energy, to transform into products. When a catalyst is present, it creates a new pathway with a lower energy barrier. This means that a greater proportion of the reactant particles will have enough energy to react, leading to an increase in the rate of reaction.

Catalysts can be either homogeneous or heterogeneous. Homogeneous catalysts are in the same phase (solid, liquid, or gas) as the reactants, while heterogeneous catalysts are in a different phase. The catalysts work by forming temporary bonds with the reactant molecules, which helps to orient them in a way that makes the reaction more likely to occur.

It's important to note that while catalysts speed up reactions, they do not affect the position of the equilibrium in a reversible reaction. This is because they increase the rate of both the forward and reverse reactions equally. Therefore, the final proportions of reactants and products remain the same, but they are reached more quickly.

In summary, catalysts play a crucial role in increasing the rate of chemical reactions by providing an alternative reaction pathway with a lower activation energy. This fundamental concept is key to understanding many processes in chemistry, from industrial synthesis to biological metabolism.