Collision Theory

• Collision theory explains how chemical reactions occur and why reaction rates differ.
• It proposes that for a reaction to occur, molecular particles must collide.
• Not all collisions lead to a reaction. Only those with sufficient energy and the correct orientation result in a successful reaction.
• This necessary energy is known as the Activation Energy (Ea). Reactions only occur if the kinetic energy of the particles involved is equal to, or greater than the Activation Energy.
• Increasing temperature provides molecular particles with more kinetic energy, leading to more successful collisions per second, thus increasing the rate of reaction.
• Similarly, increasing the concentration or pressure also increases the frequency of collisions hence the rate of reaction. This is because there are more particles per unit volume to initiate collisions.
• Particle size can also impact the rate of reaction. The smaller the particle size, the larger the surface area for collisions to occur, hence increasing the rate.
• Catalysts speed up chemical reactions without being used up. They provide an alternative reaction pathway with a lower activation energy, hence increasing the rate of reaction.
• The rate equation expression (rate=k[A]^m[B]^n) is derived from collision theory. It describes how the rate of reaction is affected by the concentration of reactants.
• Finally, a ‘successful’ collision is one where reactant particles collide with the correct orientation and with sufficient energy to overcome the activation energy barrier. If both conditions are not satisfied, the particles will simply bounce off each other and no reaction will occur.