Collision theory and rate graphs

Last time

Before we start: last time, in Equilibrium, the Haber and Contact processes, you learned these points. Here they are again on one card. Keep them in mind, because today builds on them.

Fast or slow?

Iron rusts over years. Magnesium ribbon burns in a flash. A firework explodes in a fraction of a second. Why are some reactions fast and others slow? One idea explains all of them, and it is about collisions.

Collision theory

For a reaction to happen, the particles must collide with each other. But a collision is not enough. The particles must collide with at least the activation energy. A collision with enough energy is called a successful collision. The rate of reaction depends on how many successful collisions there are each second.

Concentration

A higher concentration means more particles in every cubic decimetre. They are closer together, so they collide more often. More collisions every second means more successful collisions every second, and a faster rate. For gases, raising the pressure does the same thing. The particles are squashed closer together.

Surface area

Only the particles at the surface of a solid can be hit by the other reactant. Cutting a lump into a powder exposes many more particles, so the surface area is larger. More particles exposed means more collisions each second, so the reaction is faster.