Collision Theory


Core Concepts-Collision Theory

In this tutorial on collision theory, you will learn about what makes a collision successful. You will also learn about the ways to increase collisions and reaction rates, as well as the importance of collision theory.

Topics Covered in Other Articles

Understanding Kinetic Molecular Theory

Reaction Rates and How to Determine Rate Law

Catalysts and Activation Energy


States of Matter

Vocabulary for Collision Theory

Catalyst – A substance that is not used up in a reaction, but decreases activation energy and speeds up a reaction. Learn more about catalysts here.

Kinetic Energy – The energy of an object due to motion

Activation Energy – The minimum energy a reactant must possess in order to react

Kinetics – The study of the rate and speed of reactions

Introduction to Collision Theory

Collision theory states that for a reaction to take place, reactants must collide properly. The rate of reaction is equal to the frequency of collisions. Collision theory is limited to gases because frequencies of atomic collisions can only be calculated accurately with gases. Max Trautz proposed collision theory in 1916, as well as William Lewis independently in 1918.

Effective Collisions

An effective collision is one that produces a chemical reaction. The higher the rate of effective collisions, the faster the reaction rate.

For a collision to be effective in producing chemical change, it must have enough energy and the correct orientation. Successful collisions must have enough kinetic energy to disrupt and rearrange the bonds between atoms or they will simply bounce off of each other like balls. The frequency of collisions between two reactants is proportional to the concentration of the reactants. For example, doubling the concentration of one reactant will double the number of collisions. At room temperature, one cubic centimeter of gas has 1033 collisions a second. However, if all of these collisions were successful, all reactions would be complete within a second.

Orientation of collisions becomes more important the more complex the molecules in the reaction are. For example, in the reaction between N2O and NO, the oxygen of N2O must hit the nitrogen of NO for a reaction to happen. The more complex the molecules, the fewer number of collisions will have the proper orientation for a reaction. 

In summary, for a collision to be effective to cause a reaction:

  • Molecules must collide
  • Collisions must have enough energy (activation energy)
  • Molecules must have the proper orientation

Increasing Collisions

There are many factors that can be manipulated to increase the number of collisions, and therefore the rate of the reaction. Some of these factors are listed and explained below.

Concentration of Reactants

The higher the concentration of reactants, the more molecules available to collide.



Higher temperatures cause reactants to have more kinetic energy. With more kinetic energy, molecules will move faster. Therefore, reactants will collide more often and with more energy.


By increasing the pressure, the space between molecules of reactants decreases. With less space, molecules will collide more frequently.

States of Matter

Gasses are quicker than liquids, which are quicker than solids. For example, if people are running in random directions in a room, they are more likely to collide than if they are walking.


Catalysts speed up reactions and change the mechanism by which molecules collide. They are not consumed by the reaction and remain unchanged. Catalysts speed up the reaction by lowering the activation energy required for a reaction.

The energy of a reaction with and without a catalyst present to lower activation energy

Why is Collision Theory Important?

Collision theory explains how to increase reaction rates. By analyzing all of the factors that increase collisions, mathematical equations are used to form rate equations. Reaction rates are dependent on the factors that increase collisions. It is important to be able to determine and manipulate the rate of a reaction to effectively create needed products.