How do catalysts work?

How do catalysts work?

A catalyst increases reaction rates in a slightly different way from other methods of increasing reaction rate. The function of a catalyst is to lower the activation energy so that a greater proportion of the particles have enough energy to react. A catalyst can lower the activation energy for a reaction by:

• orienting the reacting particles in such a way that successful collisions are more likely

• reacting with the reactants to form an intermediate that requires lower energy to form the product

Some metals e.g. platinum, copper and iron can act as catalysts in certain reactions. In our own bodies, we have enzymes that are catalysts, which help to speed up biological reactions. Catalysts generally react with one or more of the reactants to form a chemical intermediate, which then reacts to form the final product. The chemical intermediate is sometimes called the activated complex.

The following is an example of how a reaction involving a catalyst might proceed. A and B are reactants, \(\color{blue}{\text{C}}\) is the catalyst, and D is the product of the reaction of A and B.

Step 1: \({\text{A}} + \color{blue}{\text{C}} \to \text{A}\color{blue}{\text{C}}\)

Step 2: \(\text{B} + \text{A}\color{blue}{\text{C}} \to \text{A}\color{blue}{\text{C}}\text{B}\)

Step 3: \(\text{A}\color{blue}{\text{C}}\text{B} \to \color{blue}{\text{C}} + \text{D}\)

\(\text{A}\color{blue}{\text{C}}\text{B}\) represents the intermediate chemical. Although the catalyst (\(\color{blue}{\text{C}}\)) is consumed by reaction 1, it is later released again by reaction 3, so that the overall reaction with a catalyst is as follows:

\(\text{A} + \text{B} + \color{blue}{\text{C}} \to \text{D} + \color{blue}{\text{C}}\)

You can see from this that the catalyst is released at the end of the reaction, completely unchanged. Without a catalyst the overall reaction would be:

\(\text{A} + \text{B}\) \(\to\) \(\text{D}\)

The catalyst has provided an alternative set of reaction steps, which we refer to as an alternative pathway. The pathway involving the catalyst requires less activation energy and is therefore faster.

This can be seen in the following diagram (see figure below).

Energy diagrams are useful to illustrate the effect of a catalyst on reaction rates. Catalysts decrease the activation energy required for a reaction to proceed (shown by the smaller magnitude of the activation energy on the energy diagram in the figure below), and therefore increase the reaction rate. Remember that with a catalyst, the average kinetic energy of the molecules remains the same but the required energy decreases (see image above).