Catalytic action and activation energy

For a reaction to occur, colliding particles need enough energy to overcome an energy barrier. The minimum energy needed is called the activation energy.

A catalyst speeds up a reaction by providing a different pathway from the uncatalysed reaction. This alternative pathway has a lower activation energy. At the same temperature, a greater proportion of particle collisions have enough energy for the reaction to occur, so more successful collisions happen each second and the reaction is faster.

Catalysts on a reaction profile

A reaction profile shows how energy changes along a reaction pathway. The rise from the reactants’ energy level to the peak represents the activation energy. The catalysed pathway has a lower peak than the uncatalysed pathway, showing its lower activation energy.

Both pathways start with the same reactants and end with the same products. The catalyst changes the pathway, not the energy levels of the reactants or products. Therefore, the overall energy change between reactants and products stays the same.

Catalysts

A catalyst is a substance that changes the rate of a chemical reaction, usually making it faster. It can take part in the reaction, but it is not used up overall: it remains chemically unchanged at the end.

A substance can therefore be identified as a catalyst if it increases the reaction rate while remaining chemically unchanged at the end. Because it is not consumed or produced overall, a catalyst is not written as a reactant or product in the balanced chemical equation.

Catalyst specificity

Catalysts are specific: a catalyst that speeds up one reaction may not work for another. Different reactions therefore require different catalysts.

In living organisms, enzymes are catalysts that speed up biological chemical reactions.