How does collision theory explain reaction rates?

Collision theory explains reaction rates by stating that reactions occur when particles collide with sufficient energy and correct orientation.

In more detail, collision theory is a fundamental concept in chemistry that helps us understand how reactions occur and why reaction rates differ. According to this theory, for a reaction to take place, the reactant particles must collide. However, not all collisions cause a reaction. The particles must collide with a certain minimum energy, known as the activation energy, and in the correct orientation for a successful reaction to occur.

The rate of a chemical reaction is directly proportional to the number of successful collisions per second. Therefore, anything that increases the number of successful collisions will increase the rate of reaction. This can include increasing the concentration of reactants, raising the temperature, increasing the surface area of solid reactants, or adding a catalyst.

Increasing the concentration of reactants or the pressure in a system of gaseous reactants will increase the number of particles in a given volume. This leads to more collisions and therefore a higher reaction rate. Similarly, increasing the temperature provides particles with more kinetic energy, making collisions more frequent and energetic, thus increasing the reaction rate. Increasing the surface area of a solid reactant exposes more particles to collide with reactant particles in the solution, again increasing the reaction rate. Finally, a catalyst works by providing an alternative reaction pathway with a lower activation energy, allowing more collisions to result in a reaction and thus increasing the reaction rate.

In summary, collision theory provides a framework for understanding reaction rates by considering the frequency, energy and orientation of collisions between reactant particles.