Collision Theory & What Controls Rate

A chemical reaction only happens when particles collide, with enough energy (the activation energy), and in the correct orientation. Anything that increases the number of successful collisions per second will speed up the reaction.

Raising the temperature gives particles more kinetic energy. They move faster → collide more often, and a larger proportion of collisions exceed the activation energy. Both effects increase the rate.

Higher concentration packs more reactant particles into the same volume. Smaller solid pieces expose more surface. In both cases: more collisions per second → faster reaction.

A catalyst provides an alternative pathway with a lower activation energy. It is not consumed, and it does not increase particle energy — it just makes more collisions successful.

The gradient = the rate. The reaction is fastest at the start (steepest part) and slows as reactants are used up. A flat line means the reaction has finished.