Collision Theory (Cambridge O Level Chemistry)

• Collision theory states that in order for a reaction to occur:
  • The particles must collide with each other
  • The collision must have sufficient energy to cause a reaction i.e. enough energy to break bonds
• The minimum energy that colliding particles must have to react is known as the activation energy
• Collisions which result in a reaction are known as successful collisions
  • If they have sufficient energy (i.e. energy greater than the activation energy), they will react, and the collision will be successful
• Not all collisions result in a chemical reaction:
  • Most collisions just result in the colliding particles bouncing off each other
  • Collisions which do not result in a reaction are known as unsuccessful collisions
• Unsuccessful collisions happen when the colliding species do not have enough energy to break the necessary bonds (i.e. they collide with energy less than the activation energy)

Diagram showing a successful and an unsuccessful collision

• Increasing the number of successful collisions means that a greater proportion of reactant particles collide to form product molecules
• The number of successful collisions depends on:
• 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 - fewer collisions will have an energy that exceeds higher activation energy and fewer collisions will be successful
• These all have an impact on the rate of reaction which is dependent on the number of successful collisions per unit of time