MCQ Detailed View

In chemical kinetics, a reaction mechanism is the step-by-step sequence of elementary reactions that describes how a chemical reaction occurs. During this process, different species may appear temporarily and then disappear before the overall reaction is complete. These temporary species are called reaction intermediates.

An intermediate is defined as a chemical species that is produced in one step of a reaction mechanism and then consumed in a later step. Because they are used up before the reaction ends, intermediates never appear in the overall balanced chemical equation. They are often unstable and cannot usually be isolated under normal reaction conditions. Common examples include free radicals, carbocations, carbanions, and complex ion species.

It is important to distinguish intermediates from other terms in reaction kinetics. A catalyst also participates in reactions but is regenerated at the end of the mechanism, so it appears unchanged in the overall reaction. An inhibitor slows down or prevents a reaction but does not act as a temporary product like an intermediate. The activated complex, also called the transition state, is a high-energy arrangement of atoms at the top of the energy barrier, which exists only momentarily and cannot be isolated.

For example, in the two-step reaction:

1. 
   

   NO₂ + F₂ → NO₂F + F

   
   


2. 
   

   F + NO₂ → NO₂F

   
   

The fluorine atom (F) is produced in step one and consumed in step two. It does not appear in the overall equation: 2NO₂ + F₂ → 2NO₂F. Here, fluorine (F) is the intermediate.

Understanding intermediates is essential in physical chemistry, particularly in reaction kinetics, catalysis, and mechanism studies, because they explain how complex reactions proceed through multiple steps.