collision theory

Higher and foundation tiers

Collision theory- explaining the rates of chemical reactions

Before a chemical reaction can happen between two reacting substances the particles present in each of them must collide with a large enough force to break all the bonds holding the reactant particles together. As an example consider the reaction between hydrogen and chlorine to make hydrogen chloride gas. The image below shows a molecule of hydrogen and a molecule of chlorine reacting to make two new molecules of hydrogen chloride gas.

Before a chemical reaction happens the reacting particles must collide with enough force to break the bonds holding the molecules together.

Before the hydrogen and chlorine gases can react you need to think about what must happen during the reaction:

A successful collision is one in which the chemical bonds in the reactant molecules are broken.  The molecules collide with energy above the activation energy.

If the molecules collide with less force and the bonds in the reactants do not break, this is called an unsuccessful collision.

Particle picture for an unsuccessful collision

High activation energies will effectively stop the reactants turning into products. High activation energies will likely result in very slow reactions while fast reactions; such as explosions are likely to have low activation energies. Any factor that increases the number of successful collisions in a reaction will increase the rate of that reaction. There are four factors that will affect the rates of chemical reactions; these are:

Practice questions

Check your understanding - Questions on collision theory and reaction rates

Check your understanding - Quick quiz on rates and collision theory.

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