Intermediates and transition states play crucial roles in the mechanisms of chemical reactions.

Intermediates are stable molecules that are formed during a reaction. They can be isolated and studied, exhibiting a longer lifetime compared to transition states. Intermediates arise when reactants collide and create a new molecule that is not the final product. This intermediate can further react to yield the ultimate product. For instance, in the reaction between hydrogen and oxygen to produce water, the hydroxyl radical ($\text{OH}$) serves as the intermediate, generated when hydrogen and oxygen interact.

In contrast, transition states are transient and unstable entities that occur during a reaction. They represent the highest energy point in the reaction pathway, embodying the state of maximum energy. Transition states form when reactants collide, resulting in a new molecule that is inherently unstable and quickly decomposes to form the final product. The significance of transition states lies in their influence on the reaction rate; specifically, the lower the energy of the transition state, the faster the reaction proceeds. For example, in the hydrogen-oxygen reaction to form water, the transition state is known as the activated complex, which forms when the reactants collide and create an unstable molecular configuration that swiftly converts into water.

In summary, intermediates are stable molecules produced during a reaction, while transition states are unstable, short-lived molecules that mark the peak energy point of the reaction. Understanding both intermediates and transition states is essential for comprehending the dynamics of chemical reactions.