Introduction to Reaction Mechanisms

Understanding Reaction Mechanisms

A reaction mechanism refers to the step-by-step sequence of elementary reactions by which the overall chemical change occurs.
Elementary reactions are processes in which reactants are converted into products in a single step with a single transition state.
Reaction mechanisms describe the molecular movements and the sequence of chemical reactions occurring on the atomic level.

Reactions may occur in one step or multiple steps, the latter is known as a multistep reaction, which consists of a number of elementary reactions.
Each elementary step will have its own activation energy and transition state.
Activation energy is the minimum amount of energy required for a reaction to proceed.
The transition state in a chemical reaction is a high-energy state that must be achieved for the reaction to proceed.

An intermediate is a species that is formed during the reaction but gets consumed before the reaction is completed.
Intermediates are very essential in multistep reactions; they allow the reaction to proceed in steps to reach the final product(s).
Even though they do not appear in the net reaction, they play a crucial role in the overall reaction mechanism.

The rate determining step is the slowest step in the reaction and it determines the speed at which the entire reaction occurs.
The rate determining step has the highest activation energy in a multistep reaction.
Means the rate of the overall reaction is determined by the rate of the slowest elementary step in the reaction mechanism.

The molecularity of a reaction step is determined by how many molecules are involved in that step.
Each elementary step in a mechanism can be described as being unimolecular, bimolecular, or termolecular.
Unimolecular steps involve a single molecule, bimolecular steps involve two molecules, and termolecular steps involve three molecules. Termolecular reactions are relatively rare because of the low probability of three molecules colliding simultaneously.