With very few exceptions, reaction rates increase with increasing temperature. For example, the time required to half – boil an egg in water is much shorter if the "reaction" is carried out at 100°C (about 10 min) than at 80°C (about 30 min). Conversely, an effective way to preserve foods is to store them at subzero temperatures, thereby slowing the rate of bacterial decay.
The rate constants can be explained by three theories. They are:
- Collision theory
- Arrhenius theory
- Transition state theory
The main postulates of collision theory
The main postulates are:
- Reacting molecules shall have to collide for any reaction to occur.
- All collisions do not lead to the formation of the products.
- The colliding molecules shall have to possess a minimum energy to give products. This minimum energy is called threshold energy. This is higher than that of the molecules in the Normal state.
- The energy of the molecules at STP is very much less than this Threshold energy.
- The difference between the Threshold energy and the energy of molecules in the Normal state is Activation energy.
Activation energy = [Threshold energy – energy of the normal molecules] - The molecules possessing the threshold energy are called activated molecules. These are formed in small numbers during collisions between normal molecules.
- Collisions occurring between activated molecules are called activated collisions. Activated collisions alone lead to the formation of the products in the reaction.
- The reaction of the activated collisions amongst the total collisions is very much small.
Collision theory of Chemical Kinetics:
Collision theory is a theory proposed by William Lewis in 1916 and 1918, that qualitatively explains how chemical reactions occur
and why reaction rates differ for different reactions. This theory postulates that gas molecules frequently collide with one
another. Therefore, it seems logical to assume – and it is generally true – that chemical reactions occur as a
result of collisions between reacting molecules. In terms of the collision theory of chemical kinetics, then, we expect the
rate of a reaction to be directly proportional to the number of molecular collisions: rate is proportional to number of collisions
per second.
In order to take place a reaction, the molecules involving in the reaction must collide with each other. These molecules must possess a total kinetic energy equal to or greater than the activation energy (Ea), which is the minimum amount of energy required to initiate a chemical reaction. When molecules collide they form an activated complex (also called the transition state), a temperature species formed by the reactant molecules as a result of collision, before they form the product.
The below figure shows two different potential energy profiles for the reaction,
A + 
AB
C + D
AB denotes an activated complex, formed by the collision between A and B. If the products are more stable than
the reactants, then the reaction will be accompanied by release of heat i.e., the reaction is exothermic. On the other hand,
if the products are less stable than the reactants, then heat will be absorbed by the reacting mixtures from the surroundings.
In both the cases, we plot the potential energy of the reacting system versus the progress of the reactions. Qualitatively,
these plots show the potential energy changes as reactants are converted to products.
Activation energy is a barrier which prevents less energetic molecules from reacting. Normally, only a small fraction of the
colliding molecules i.e., the fast moving ones have enough kinetic energy to exceed the activation energy.