Catalysis

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Catalysis

A catalyst speeds up a chemical reaction without itself undergoing any permanent chemical change. They are usually specific to one particular reaction and they speed up both forward and back reactions equally.

All types of catalyst work by providing an alternative pathway that has a lower activation energy than the original one. The catalyst becomes involved in the reaction mechanism but is reformed at the end.

Consider the reaction: A + B → AB

Catalysis

In the presence of a catalyst a new pathway is created.

A + Catalyst → A-Catalyst

A-Catalyst + B → A-B + A-Catalyst

The new steps involve two transition states each with a much lower activation energy than the uncatalysed transition state. Hence, at a particular temperature, far more molecules will have sufficient energy to follow the catalysed pathway and so the rate will be faster.

Homogenous catalysts function in the same state as the reactants (usually all in aqueous solution).

Example:

CH3COOH(aq) + C2H5OH(aq) → CH3COOC2H5(aq) + H2O(l) is catalysed by the acid: H+(aq)

Enzymes are important homogenous catalysts.

Heterogenous catalyst function in a different state to the reactants. Many transition metals act as solid catalysts in solution or gas reactions.

This type of catalyst functions by absorbing the reactants onto its surface and forming weak bonds with them. This has the effect of bringing the reactant molecules closer together, weakening their bonds and thus lowering the activation energy.

Note: Catalysts never appear in the stoichiometric equation. However, they may appear in the rate equation if they take part in the rate determining step.

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