Collision Theory & Activation Energy (AQA GCSE Chemistry)

Collision theory

• Particle theory states that chemical reactions occur only when the reactant particles collide with sufficient energy to react

  • The minimum amount of energy needed is called the activation energy

  • Activation energy is different for each reaction

• Collisions can be described as successful or unsuccessful

A successful collision

• A successful collision means that the reactant particles colliding have sufficient energy, i.e. greater than or equal to the activation energy

• This means that the reactant particles rearrange to form the products

The collision is successful resulting in a rearrangement of atoms to form the products

An unsuccessful collision

• An unsuccessful collision means that the reactant particles have insufficient energy, i.e. less than the activation energy

• This means that the reactant particles just bounce off each other and remain unchanged

The collision is unsuccessful resulting in no rearrangement of atoms  

• Increasing the number of successful collisions means that a greater proportion of reactant particles collide to form product molecules

• The following all affect the rate of reaction which is dependent on the number of successful collisions per unit time:

  • The number of particles per unit volume - more particles in a given volume will produce more frequent successful collisions

  • The frequency of collisions - a greater number of collisions per second will give a greater number of successful collisions per second

  • The kinetic energy of the particles - greater kinetic energy means a greater proportion of collisions will have an energy that exceeds the activation energy and the more frequent the collisions will be as the particles are moving quicker, therefore, more collisions will be successful

  • The activation energy - if a reaction has a high activation energy, there will be fewer collisions with an energy that exceeds the activation energy and fewer collisions will be successful

• So, the rate of a reaction is dependent on the energy of collisions as well as the number of collisions

Explaining rates

• Increasing the number of successful collisions means that a greater proportion of reactant particles collide to form product molecules

• We have seen previously that the following factors influence the rate of reaction

  • Increasing concentration / pressure

  • Increasing temperature

  • Increase the surface area of a solid reactant

  • Use of a catalyst

• We can use collision theory to explain why these factors influence the reaction rate:

How increasing concentration affects rate

• Increasing the concentration of a solution increases the rate of reaction 

• Increasing the concentration means that there are more reactant particles in a given volume

  • This causes more collisions per second

  • Leading to more frequent and successful collisions per second

  • Therefore, the rate of reaction increases 

• If you double the number of particles, you will double the number of collisions per second

  • The number of collisions is proportional to the number of particles present

Diagram showing the effect of increasing concentration

A higher concentration of particles in (b) means that there are more particles present in the same volume than (a) so the number of collisions and successful collisions between particles increases causing an increased rate of reaction

How increasing pressure affects rate

• Increasing the pressure of a gas increases the rate of reaction 

• Increasing the pressure means that there are the same number of reactant particles in a smaller volume

  • This causes more collisions per second

  • Leading to more frequent and successful collisions per second

  • Therefore, the rate of reaction increases 

Diagram showing the effect of increasing pressure

The higher pressure (b) means that there are the same number of particles present in a smaller volume than (a) so the number of collisions and successful collisions between particles increases causing an increased rate of reaction

How increasing temperature affects rate

• Increasing the temperature increases the rate of reaction 

• Increasing the temperature means that the particles have more kinetic energy

  • This causes more collisions per second

  • Leading to more frequent and successful collisions per second

  • Therefore, the rate of reaction increases

• The effect of temperature on collisions is not so straightforward as concentration or surface area; a small increase in temperature causes a large increase in rate

• For aqueous and gaseous systems, a rough rule of thumb is that for every 10 ^oC increase in temperature, the rate of reaction approximately doubles

Diagram showing the effect of increasing temperature

An increase in temperature causes an increase in the kinetic energy of the particles. The number of successful collisions increases 

How increasing the surface area affects rate

• Increasing the surface area increases the rate of reaction 

• Increasing the surface area means that a greater surface area of particles will be exposed to the other reactant

  • This causes more collisions per second

  • Leading to more frequent and successful collisions per second

  • Therefore, the rate of reaction increases

• If you double the surface area, you will double the number of collisions per second

Diagram showing the effect of increasing surface area

An increase in surface area means more collisions per second