Collision theory

Collision theory, theory used to determine the rates of chemical reactions, particularly for gases. The collision theory is based on the assumption that for a reaction to occur, it is necessary for the reactants (atoms or molecules) to come together or collide with one another. Not all collisions, however, bring about chemical change. Only a small fraction of the total number of collision is effective.

Collision theory states that for a chemical reaction to occur, the reacting particles must collide with one another. The fraction of the collision that results in a reaction is called successful collision.

The successful collisions must have enough energy, also known as activation energy, at the moment of impact to break the pre-existing bonds and form all new bonds. This results in the formation of product(new substance).

  Increasing the concentration of the reactant brings about more collisions and hence more successful collisions. Increasing the temperature increases the average kinetic energy of the molecules in a solution, increasing the amount of collisions that have enough energy. Collision theory was proposed independently by Max Trautz in 1916 and William Lewis in 1918.

Activation energy can be defined as the minimum amount of energy possessed by the colliding particles for a chemical reaction to occur. It is equivalent (equal to) the energy barrier that must be broken before bonds are stretched and successfully broken to enable the reaction take place.