A reaction can occur when molecules of reactants collide with each other to form an unstable intermediate (Fig. 20.12). The intermediate exists for a very short time and then breaks up to form product molecules. The energy required to form this intermediate, called activated complex (C), is known as activation energy (E a). The distribution of kinetic energy of different fraction (NE/NT) of molecules is shown in Fig. 20.12.
Fig. 20.12. Distribution curve showing energies among gaseous molecules.
Only those collisions result in the formation of products which possess energy equal to or more than the certain minimum energy called threshold energy. Collisions of the molecules possessing energy less than threshold energy do not form products. It means that between reactants and products there is an energy barrier which must be crossed before the reactants are converted into products. The energy required for crossing this energy barrier is supplied by the kinetic energy of the -molecules.
“The minimum extra energy over and above the average potential energy of the reactants which must be supplied to the reactants to enable them to cross over the energy barrier between reactants and products is called Activation energy”. Thus,
Activation energy= (Threshold energy) – (Average energy of the reactants)
or E a = E y-E R
The idea of activation energy and the energy barrier involved in a reaction is given in the Fig. 20.13.
Fig. 20.13. Illustration of activation energy and energy barrier involved in a reaction.
It is important to note that each reaction has a definite value of Ea and this decides the fraction of total collisions which are effective. Obviously, if the activation energy for a reaction is low, large number of molecules can have this energy and the fraction of effective collision, J, will be large. Such a reaction proceeds at high rate. On the other hand, if the activation energy is high, then f will be small and the reaction may be quite slow.
For fast reactions; activation energies are low.
For slow reactions; activation energies are high.
For example, for the reaction between NO and O2, Ea is low and hence, reaction is fast whereas for the reaction between CO and O2, Ea is high and hence, the reaction is slow.
2NO + O2 à 2NO2
2CO + O2 à 2CO2
The two theories can be successfully used to explain how the factors of surface area, concentration or pressure of gases, temperature and catalysts affect reaction rates.
In the case of surface area, concentration and pressure of gases at constant temperature, any increase in any individual value is equivalent to an increase in the number of molecules colliding per second. An increase in the frequency of effective collisions results in an increase in reaction rate.
ACTIVATED COMPLEX OR TRANSITION STATE THEORY
It has been pointed out earlier that during the chemical reaction certain bonds are broken and certain new bonds are formed. The breaking of bonds requires energy whereas the formation of bonds results in the release of energy. For example, in the reaction of hydrogen with iodine to form hydrogen iodine, when a molecule of hydrogen approaches that of iodine, H-H and I- I bonds start breaking and H- I bonds start forming. In the beginning, breaking of bonds predominates and therefore, energy of the system starts increasing till it reaches a maxima (corresponding to threshold energy). After this, the energy starts decreasing because the process of bond· formation predominates and finally leads to the product hydrogen iodide. The arrangement of atoms corresponding to energy 11UlXima (threshold energy) is called transition state or activated complex. In transition state, the system has partial reactant character and partial product character as shown in Fig. 20.14 and Fig. 20.15.
Fig. 20.14. Formation of activated complex during the reaction of H2 and I2 from HI.
Fig. 20.15. Transition state or activated complex.
The difference between energy of the transition state and energy of the reactants is equal to activation energy.
E transition state - E reactants = E activation