Catalysts speed up chemical reactions by lowering activation energy, but not all catalysts work the same way. In IB Chemistry, catalysis appears in kinetics (Topic 6), but you also see it in industrial processes, environmental chemistry, and organic mechanisms. Understanding the difference between heterogeneous and homogeneous catalysis helps you explain reaction pathways and evaluate industrial methods in exams.
What Is a Catalyst?
A catalyst is a substance that increases the rate of reaction without being consumed.
It provides an alternative pathway with a lower activation energy, increasing the number of successful collisions.
Catalysts fall into two major categories:
- Heterogeneous catalysts
- Homogeneous catalysts
Their names describe their physical state relative to the reactants.
Heterogeneous Catalysis
Definition
A heterogeneous catalyst is in a different physical state from the reactants.
Example:
Reactants = gases
Catalyst = solid metal surface
This is very common in industrial chemistry.
How It Works
Heterogeneous catalysis typically occurs on a solid surface. The mechanism involves:
- Adsorption – Reactant molecules stick to the catalyst surface.
- Activation – Bonds weaken due to the surface interaction.
- Reaction – Reactants combine more easily.
- Desorption – Products leave the surface, freeing active sites.
Because reactants interact directly with the surface, the catalyst must have a large surface area.
