1. Catalysts speed up reactions by lowering the activation energy required for a reaction to occur.
2. They do this by giving an alternative route (reaction pathway) that requires less energy.

The activation energy would be lower in the presence of a catalyst.

A catalyst that has the same physical state as the reactants.

For example, sulfuric acid is a homogeneous catalyst in the production of esters because sulfuric acid and the reactants (an alcohol and a carboxylic acid) are all liquids.

A catalyst that has a different physical state from the reactants.

For example, iron is a heterogeneous catalyst in the Haber process because iron is a solid and the reactants (N₂ and H₂) are gases.

The economic benefits of using catalysts are:

1. They increase the rate of reactions, allowing more products to be manufactured and sold.
2. They reduce energy requirements by enabling reactions to occur at lower temperatures and pressures.
3. They improve atom economy, which minimizes material wastage, leading to cost savings.

Using catalysts can improve sustainability by:

1. Enabling reactions at lower temperatures, which results in reduced emissions of greenhouse gases, such as CO₂.
2. Improving atom economy, which cuts down on the waste produced during chemical processes.
3. Accelerating the breakdown of pollutants from industrial processes, reducing their environmental impact.

When a catayst is added, the shape of the Boltzmann distribution curve remains the same. The activiation energy is lowered.