Multi-Step Reactions

A balanced equation shows only reactants and products, but a reaction mechanism reveals the step-by-step molecular process. The mechanism consists of elementary reactions that occur in sequence and sum to yield the overall equation. This detailed view explains how bonds break and form at each stage, providing insight into the actual chemical transformation.

Reaction intermediates are molecules formed in one elementary step and consumed in another, appearing in the mechanism but not in the overall equation. Transition states, by contrast, exist only momentarily during the transformation of reactants to products and represent the highest-energy point. Both are temporary species, but intermediates persist between steps while transition states are instantaneous.

The rate-determining step is the slowest elementary step in a multi-step mechanism and controls the overall reaction rate. A reaction cannot proceed faster than its slowest step, so this step limits the net rate. The rate law derived from the rate-determining step can be compared to experimental data to verify whether a proposed reaction mechanism is correct.

Derive the rate law from the rate-determining step and compare it to the experimentally determined rate law. If they match, the mechanism is consistent with experimental observations. The presence of reaction intermediates in the mechanism but not in the overall equation further supports the validity of the proposed mechanism.

A second-order rate law indicates the reaction rate depends on the square of one reactant's concentration. This typically suggests that either two molecules of that reactant collide in an elementary step, or the rate-determining step involves a reactant whose concentration is affected by a prior equilibrium. This relationship cannot be predicted from the overall balanced equation alone.

When a reversible elementary step precedes the rate-determining step, the rate law becomes more complex because reaction intermediates appear in the expression. The concept of equilibrium—where forward and reverse reaction rates are equal—helps derive the overall rate law by relating intermediate concentrations to reactant concentrations.

Balanced equations represent only the net chemical change and do not reveal the actual molecular pathway. Rate laws derived directly from balanced equations often do not match experimental data, indicating a multi-step mechanism. Experimental rate laws must be measured first, then used to deduce the correct reaction mechanism and elementary steps.