Collision theory

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Kinetic-Molecular theory can be used to calculate the collision frequency. In a gas at standard conditions, in 1 cm3 there are about 1030 collisions per second or each molecule under goes about 1010 collisions/s. If each collision produced a reaction, the rate would be about 106 M s-1. Actual fast rates are on the order of 104 M s-1. Still a very rapid rate, but meaning that only about 1% of the collisions result in products being formed. For slower reactions and even smaller percentage of collisions between reactant leads to products.

Collision theory explains why most collisions do not result in a reaction.

Collision theory is based on three postulates:

  1. Chemical reactions in the gas phase are due to the collision of the reactant particles.
  2. A collision results in a reaction only if certain threshold energy is exceeded.
  3. A collision results in a reaction only if the colliding particles are correctly oriented to one another.

In 1888, Swedish chemist Arrhenius suggested that molecules must possess a certain minimum amount of energy in order to react. This energy is the kinetic energies of the colliding molecules. Upon collision, the kinetic energy of the molecules can be used to stretch, bend, and break bonds, leading to chemical reactions.

If molecules are moving slowly with little kinetic energy, they will bounce off one another without changing. In order to react, colliding molecules must have a total kinetic energy equal to or greater than some minimum value. The minimum energy required to initiate a chemical reaction is called the activation energy, Ea. The value Ea varies from reaction to reaction. Only faster moving particles will collide with sufficient kinetic energy to overcome the activation barrier.

If activation energy is high, only a few molecules have sufficient kinetic energy to overcome the barrier and the reaction is slower. As temperature increases, more of the molecules will have sufficient kinetic energy to overcome the barrier, therefore, the reaction rate increases.

At the moment of effective collision, both bond breaking and bond formation are occurring. As bond rearrangement occurs, an unstable intermediate species, called an activated complex or transition state, exists in the reaction mixture. The activation complex represents the highest energy point along the pathway from reactants to products. Because it is highly unstable, the activated complex may have a lifetime as short as 10-15 s.

However, reaction rates are smaller than the rate of collisions with enough energy. This is because the orientation of molecules is important. Reactants must be orientated to allow the formation of any new bonds needed to produce products.