Many chemical reactions, particularly those involving several overall steps, do not occur in a single, simultaneous collision of all reactant molecules. Instead, they proceed through a sequence of individual, simpler molecular events known as elementary steps. The complete sequence of these elementary steps that collectively describes the pathway from reactants to products is called the reaction mechanism. The sum of all the elementary steps in a valid mechanism must always yield the overall balanced chemical equation for the reaction.

• Elementary Steps: Each elementary step represents a single molecular event, such as a molecule breaking apart (unimolecular), two molecules colliding (bimolecular), or very rarely, three molecules colliding simultaneously (termolecular). For an elementary step, the reaction orders do correspond directly to the stoichiometric coefficients of the reactants involved in that specific step.

• Intermediates: Within a reaction mechanism, some species may be formed in one elementary step and then consumed in a subsequent elementary step. These transient species are called intermediates. They do not appear in the overall balanced chemical equation because they are neither initial reactants nor final products. Intermediates are often highly reactive and thus exist for only a very short time.

• Molecularity: The molecularity of an elementary step refers to the number of reactant molecules or ions that participate in that particular step. * Unimolecular: A single reactant molecule rearranges or decomposes (e.g., A → products). * Bimolecular: Two reactant molecules collide (e.g., A + B → products, or 2A → products). * Termolecular: Three reactant molecules collide simultaneously. Termolecular steps are extremely rare due to the very low probability of three particles colliding at the exact same time with the correct energy and orientation.

The Rate Determining Step (RDS): The Bottleneck of the Reaction In a multi-step reaction mechanism, one elementary step will invariably be significantly slower than all the others. This slowest elementary step is called the rate-determining step (RDS), or sometimes the rate-limiting step. Much like the weakest link in a chain determines the overall strength of the chain, or the slowest worker on an assembly line determines the output rate, the rate of the overall reaction is limited and effectively determined by the speed of its slowest elementary step.

Connecting Mechanism to the Rate Expression: The experimentally determined rate expression (rate law) for an overall reaction provides crucial insights into its reaction mechanism. The rate expression is derived directly from the rate-determining step. * The concentrations of the reactants involved in the rate-determining step (and sometimes concentrations of reactants from preceding fast equilibrium steps that produce an intermediate) will appear in the overall rate expression. * Intermediates will generally not appear in the overall rate expression, as they are not the initial reactants of the overall reaction. If an intermediate is involved in the RDS, its concentration must be expressed in terms of the initial reactants using an equilibrium expression from a preceding fast step.

• Example: Let's consider a hypothetical reaction: 2NO2(g) + F2(g) → 2NO2F(g). A proposed two-step mechanism for this reaction might be: Step 1: NO2(g) + F2(g) → NO2F(g) + F(g) (This is the slow, rate-determining step) Step 2: NO2(g) + F(g) → NO2F(g) (This is a fast step) Let's analyze this proposed mechanism: * Overall Reaction: If we add the two elementary steps, the intermediate species F(g) cancels out: (NO2(g) + F2(g)) + (NO2(g) + F(g)) → (NO2F(g) + F(g)) + NO2F(g) Summing and simplifying yields: 2NO2(g) + F2(g) → 2NO2F(g), which matches the overall balanced equation. * Rate Expression: Since Step 1 is identified as the slow, rate-determining step, the overall rate of the reaction is governed by the rate of this step. For an elementary step, the rate law is based on its stoichiometry. Therefore, the rate expression for the overall reaction would be: Rate = k [NO2] [F2]