The collision theory states that gaseous state chemical reactions takes place when two gas molecules smash together with enough kinetic energy. The smallest amount of energy that is needed for a thriving collision, which leads to a successful reaction, is referred to as the activation energy. Consequently, only a fraction of collisions result to successful reactions.
The collision theory is derived from the kinetic theory of gases. Thus we are only dealing with gaseous chemical reactions. Because of this, the ideal gas assumptions are applied. Moreover, we are as well making assumptions that:
- All the reacting molecules are moving through space in a straight line.
All molecules are inflexible spheres.
The reactions being discussed only takes place between two molecules.
The molecules ought to collide with one another.
In the end, the collision theory of gases offers us the rate constant for bimolecular gaseous reactions; it is equal to the rate of winning collisions. The rate of successful collisions is proportional to the small part of successful collisions multiplied by the general collision frequency.
The rate or frequency at which molecules collide is referred to as the collision frequency, Z. The unit of collision frequency is collisions/ unit of time. Given a box of molecules A and B, the collision frequency between molecules A and B is given by:
To enable a successful collision to take place, the molecules of the reactant ought to collide with sufficient kinetic energy to break original bonds and be able to form fresh bonds to turn into the molecules of the product. Therefore, it is referred to as the activation energy for the reaction; it is as well normally regarded as the energy barrier.
The portion of collisions that possess sufficient energy to exceed the activation barrier is denoted as:
The tiny proportion of successful collisions is directly proportional to the temperature and inversely proportional to the activation energy of the reaction