Collision Theory of reaction rate in chemical kinetics

Collision Theory of reaction rate

According to collision theory.

(i)                  Reactants are made up of molecules.

(ii)                Molecules are always in a state of random motion and hence, go on colliding with one another. Collision frequency (Z) is the number of intermolecular collision taking place per unit volume per second at a given temperature.

(iii)               A chemical reaction takes place due to inter-molecular collisions of reaction. In case of gases, the collision frequency is very high (10²⁵ to 10³⁰ collision s^-1).

(iv)              All intermolecular collision do not bring about the reaction. Only effective collisions bring about the reactions.

(v)                An effective collision is that in the which the colliding molecules

(a)    Possess energy equal to or greater than certain minimum value of energy known as the threshold energy and

(b)   Are properly oriented.

 Thus, for an effective collision, there are two barriers- Energy barrier and orientation barrier.

Collision theory of unimolecular reaction

 Unimolecular reactions are known. For example, consider the following first order reaction:

                             2N₂O₅ (g)   ž    4NO₂ (g) + O₂ (g)

In case, two molecules must collide in order of provide necessary activation energy, a second order rate law should result. Lindeman (1922) explained this anomaly by assuming that there exists a time lag between activation and the reaction of molecules. During the time lag, the activated molecule could either react or get deactivated. For example, Consider a general first order reaction.

                   AžProduct

               A+AžA*+A(activation)

A+A*žA+A(deactivation)

A*žProducts (reaction)

In case the time lag is long, step (iii) is slow and the reaction will following first order kinetics. On the order hand, if the time lag is long, step (ii) will be slow and the reaction will follow second order.

The effect of pressure also explains Lindeman’s theory. The rate of deactivation will be more at high pressure. But at sufficiently low pressure, all the activated molecules ill react before they get deactivated. Thus, with decreasing pressure, the reaction kinetic should change from first order to second order.

For biomolecular reaction

Applying this theory to biomolecular gaseous reactions, the expression for the rate constant is written as:

           K=P. Ze^-Ea/RT

Where, Z= Pre-expotential facter, A of Arrehenius equation

It is related to the frequency of collision between the reactant molecules.

e^-Ea/RT =  Fraction of colliding molecules which have the necessary activation energy.

P= probability or the steric factor. It is related to the specific rotation of the molecule. It also tell the deviation from the calculated value. Its value ranges from 1-10-8.

To understand it more clearly, consider the following reaction?

CO(g)+NO2(g)žCO2(g)+NO(g)

The CO and NO molecules must possess sufficient energy and they are so oriented that CO molecule is able to knock off the oxygen atom of NO2 during collision.

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