1.7.1 Collision Theory

What is Collision Theory?

Collision theory explains how chemical reactions occur and why different reactions happen at different rates. According to this theory, for a reaction to occur, particles (atoms, molecules, or ions) must collide with each other. However, not all collisions result in a reaction. Two main conditions must be met for a collision to be successful and lead to a reaction:

1. Sufficient energy: The colliding particles must have enough kinetic energy to overcome the energy barrier known as the activation energy (Ea).
2. Correct orientation: The particles must be oriented properly during the collision for the bonds to break and new bonds to form.

Activation Energy (Ea)

The activation energy (Ea) is the minimum energy that reacting particles must have for a successful reaction to take place. It represents the energy barrier that must be overcome for a chemical reaction to occur.

• Reactions only occur when the particles have kinetic energy greater than or equal to the activation energy.
• If particles collide with less energy than the activation energy, they will simply bounce off each other, and no reaction will occur.

Why Most Collisions Do Not Lead to a Reaction

In any reaction mixture, most particles do not have enough energy to meet the activation energy threshold. As a result:

• Only a small proportion of particle collisions have the required energy to result in a reaction.
• The majority of collisions are unsuccessful because the energy of the colliding particles is too low, or they do not collide in the correct orientation.