Chemical reaction speeds up when the temperature is increased. This fact can be explained on the basis of collision theory. According to this theory, molecule must collide in proper orientation in order to react but all the collision will not chemical reaction. During chemical reaction, the reactant molecules must have some extra energy and should approach in proper manner.
To illustrate this point, let us consider a reaction between hydrogen and iodine. When the reaction proceeds, the existing H-H and I-I bonds gets broken and new type of bond is formed. There must be proper collision between the reacting molecules for the formation of desired product. This is called effective collision.
While sidewise collision proceeds as below and does not give desired product.
Activated complex and activation energy:
The activated complex is an unstable state which is the midway between reactant and product.It is the transition state molecule, with some bonds partially broken and other bonds partially formed. The activated complex further decomposes to give product. The reacting molecule must overcome the energy barrier before forming the product. This energy is called the activation energy.
The sum of activation energy is called threshold energy. The use of catalyst increases the rate of reaction by providing the alternate path in which the energy of the activation of the reaction is less. The graph below shows the formation of product upon using catalyst and without the use of catalyst.