Q. Explain Collision Theory?

Ans.

Why do some reactions take place faster than others? What actually happens in a chemical reaction? Consider a gaseous reaction of the type:

A₂ (g) + B₂ (g) -> 2AB(g)

For a reaction to occur, molecules of the reactants must first come into contact or collide with one another. Furthermore, existing chemical bonds must be broken. Hence, only those collisions with sufficient energy can result in a reaction. This energy barrier to a reaction is called the activation energy of the reaction. This model of how chemical reactions take place is called the collision theory.

According to the collision theory, a reaction occurs only if molecules collide. Furthermore, the colliding molecules must possess enough energy to overcome the activation energy of the reaction.

In a high energy collision, the molecules successfully collide and overcome the activation energy of the reaction; therefore, new products are formed. However, in a low energy collision, the energy of the collision is not great enough to overcome the activation energy; hence, this type of collision does not result in the formation of new products.

For example, in the illustration below, the first collision between the HI molecules is a low energy collision, and therefore does not result in the formation of new products. However, the second collision is a high energy collision, and the HI molecules collide to form the new products of iodine (I) and hydrogen (H₂).
Low energy collision High energy collision

Molecules must also have a proper orientation when they collide if a reaction is to take place. Consider this reaction:

CH₃I + Br -> I + CH₃Br

Proper orientation must is essential for this reaction to occur. If Br approaches from the CH3 side of CH3I, then the reaction is favored. However, if Br approaches from the I side, a collision is not as likely, and therefore the chance of producing CH3Br is very little.