Question

Explain collision theory of reaction rates of bimolecular gases.

Answer 1

We know that for a reaction to occur there must be effective collision between the reactants. The collision theory was developed by William Lewis and Max Trautz in 1916-18. This theory gives the energetic and mechanistic aspect of reaction.

Complete step by step answer:
This theory assumes the reactant molecules as hard spheres and postulated that reaction will take place on collision of molecules with each other. The term 'collision frequency' depicts the number of collisions in one second in one unit volume of the reacting mixture. Activation energy is another factor that affects the rate of a chemical reaction. It is the energy that must be supplied to the reactants to undergo a chemical reaction.

Let’s consider a bimolecular elementary reaction as:

${\rm{A}} + {\rm{B}} \to {\rm{Products}}$

The rate of the above reaction is expressed as:

Rate=${Z_{{\rm{AB}}}}{e^{\dfrac{ - {E_a}}{RT}}}$…… (1)

Here, ${Z_{{\rm{AB}}}}$ represents the collision frequency of reactants (A and B), ${e^{\dfrac{ - {E_a}}{RT}}}$ represents fraction of molecules possessing energy equal or greater than activation energy $\left( {{E_a}} \right)$.

Equation (1) predicts that the value of rate constants of reactions involving simple molecules or atomic species is fairly accurate but significant derivations are observed in case of complex molecules.

Note: Always remember that all collisions do not lead to the formation of products. The collision in which reactant molecules possess sufficient kinetic energy and at proper orientation facilitates the breaking of bonding between reactants and formation of bonds to form products is termed as effective collision. If there is improper orientation of the reactant molecules, then there is no product formation.