Collision theory


2. What is collision theory?

Collision theory explains why chemical reactions happen and why some reactions are faster than others.

For a reaction to occur, reacting particles must collide with:

  1. Energy equal to or greater than the activation energy.

  2. The correct orientation, where the reaction requires a particular arrangement of particles.

A collision that results in a reaction is called a successful collision.

What is activation energy?

Definition to learn:

Activation energy, E_a, is the minimum energy that colliding particles must possess for a reaction to occur.

Activation energy is measured in:

Why do most collisions not result in a reaction?

Most collisions are unsuccessful because:

  • The particles do not have sufficient energy to overcome the activation energy.

  • The particles may not collide in a suitable orientation.

Not every collision needs to be successful for a reaction to proceed. The rate depends on how frequently successful collisions occur.

The frequency of successful collisions

Higher frequency of successful collisions, higher reaction rate

This is one of the most important relationships in the topic.

3. Factors that affect the rate of reaction

The main factors you need to understand are:

  • Temperature.

  • Concentration of a solution.

  • Pressure of a gas.

  • Surface area of a solid.

  • Catalysts.

A. Increasing temperature

When temperature increases:

  1. The particles gain greater average kinetic energy.

  2. They move faster.

  3. Collisions occur more frequently.

  4. A much greater proportion of particles have energy equal to or greater than E_a.

  5. The frequency of successful collisions increases.

  6. The reaction rate increases.

AQA exam wording: The key point is that the proportion of particles with energy greater than or equal to the activation energy increases significantly.

Do not just say that particles move faster. That alone does not fully explain the increase in rate.

B. Increasing concentration

When the concentration of a solution increases:

  1. There are more reactant particles per unit volume.

  2. Particles are closer together on average.

  3. Collision frequency increases.

  4. The frequency of successful collisions increases.

  5. The reaction rate increases.

Key phrase: More frequent collisions per unit volume per unit time.

C. Increasing gas pressure

For reacting gases, increasing pressure at constant temperature:

  1. Decreases the volume available to the gas particles, if the gas is compressed.

  2. Increases the number of gas particles per unit volume.

  3. Increases collision frequency.

  4. Increases the frequency of successful collisions.

  5. Increases the reaction rate.

Important: Increasing the pressure of a gas is not the same as increasing the temperature. At constant temperature, the average kinetic energy of the gas particles remains unchanged.

D. Increasing the surface area of a solid

Breaking a solid into smaller pieces increases its exposed surface area.

For example, powdered calcium carbonate reacts faster with hydrochloric acid than the same mass of large calcium carbonate chips under otherwise identical conditions.

This happens because:

  1. More of the solid is exposed to the acid.

  2. More reactant particles can collide with the solid surface at any one time.

  3. The frequency of successful collisions increases.

  4. The reaction rate increases.

Important: The total mass of solid does not need to change. It is the exposed surface area that changes.

E. Adding a catalyst

A catalyst increases the rate of a reaction by providing an alternative reaction pathway with a lower activation energy.

This means a greater proportion of reacting particles have sufficient energy to react.

A catalyst is not consumed overall in the reaction and does not change the position of equilibrium.