Activation Energy Vs Threshold Energy

We all must have heard about "Threshold Energy" and "Activation Energy". So let us now see the difference between threshold energy and activation energy. 

From 'Collision Theory', we know that a reaction occurs when the reactant molecules collide with each other and these collisions must be an effective collision so that the reaction can occur. The collision are an effective collision if the reactant molecules must have energy more than threshold energy.

Reactant molecules have low energy at room temperature due to which their collisions are not effective and reactants do not convert into products. In this situation, we give energy (or activation energy) to reactant molecules in the form of heat, light etc. So that they absorbs this energy and their energy becomes equal to or greater than threshold energy. As a result, the reactant is easily converted into products by effective collisions. 

Thus, "Threshold energy is the summation of the average energy of the reactant molecules and the activation energy (energy absorbed by reactant molecules) as shown in the above figure". Activation energy of the reactant molecule is always equal to or lesser than the threshold energy. 

➩ T.E. = Average Energy of molecules + A.E.

     Where,

               T.E. = Represent Threshold Energy

               A.E. = Represent Activation Energy

Threshold Energy??? 

Threshold energy is the minimum amount of energy that reactant molecules must possess in order to have an effective collision.

Activation Energy??? 

Activation energy of a reaction is the difference between the energy of a transition state (or activated state) and the average energy of the reactant molecules.

Average Energy Of Reactant Molecules???

The energy acquired by the reactant molecule at the start of the reaction is called as 'Average Energy Of Reactant Molecules'. 

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