Collision theory and rates of reaction

Collision Theory

Collision theory explains how chemical reactions occur and why reaction rates differ for different reactions.
For a reaction to occur, particles must collide with a certain minimum energy, known as the activation energy.
In a collision, the original bonds break and new bonds form to produce the reaction products.
Not all collisions cause a chemical reaction. The colliding particles must have sufficient energy and correct orientation.

Increasing the concentration of reactant particles increases the collision frequency which leads to an increased reaction rate.
A higher concentration means there are more particles in the same space. This increases the number of collisions per minute.

Increasing temperature increases kinetic energy of the particles, leading to more frequent and energetic collisions.
At higher temperatures, a larger proportion of particles have energy greater than the activation energy, hence reaction rate increases.

Increasing the pressure in a gaseous system increases the number of particles within a given volume, leading to increased collision frequency and hence an increased reaction rate.

For reactions involving solids, increasing the particle’s surface area provides a greater surface for collisions to occur, increasing the reaction rate.

The Maxwell-Boltzmann distribution shows the spread of energies that particles in a gas or liquid have at a particular temperature.
The area under the curve to the right of the activation energy represents the proportion of particles with energy greater than the activation energy.
As temperature increases, not only does the number of collisions increase, but the number of successful collisions also increases, as the portion of particles with sufficient energy (activation energy or greater) increases.

A catalyst provides a different pathway for a reaction with a lower activation energy.
This means more particles have enough energy to react, increasing the rate of reaction.
Catalysts do not affect the final equilibrium position of a reaction, they just help it reach equilibrium faster. They are not consumed in the reaction.