= Reaction mechanism
{wiki}
A <reaction mechanism> describes how <reactants> become <reaction products> through specified <elementary reactions>. Adding the steps must recover the overall <chemical equation>; intermediate <chemical species> cancel. That accounting is necessary, but the proposed steps must also predict a <rate law> compatible with experiment. Several <reaction mechanisms> can pass the same test.<Chemistry 2e by Paul Flowers and others/12.6 Reaction Mechanisms>[Chemistry 2e by Paul Flowers and others; §12.6, opening discussion and Relating Reaction Mechanisms to Rate Laws: net equations, elementary steps, and experimental constraints]{ref}
For example, the stipulated steps $\mathrm A+\mathrm B\longrightarrow\mathrm I$ and $\mathrm I\longrightarrow\mathrm P$ sum to $\mathrm A+\mathrm B\longrightarrow\mathrm P$. I is a <reaction intermediate>: it is formed and then consumed. The <catalyst> C in the existing <catalytic cycle> instead starts available, participates, and is regenerated; its bound form can be a <reaction intermediate>. Both can disappear from the summed <chemical equation>, so cancellation alone does not distinguish their roles. Cancellation is algebraic bookkeeping, not an assertion that every step finishes instantaneously.
The net <chemical equation> supplies consumption ratios, not a clock. For the stipulated two-step example, directional <reaction rates> $r_1=k_1[\mathrm A][\mathrm B]$ and $r_2=k_2[\mathrm I]$ give $d[\mathrm I]/dt=r_1-r_2$ and $d[\mathrm P]/dt=r_2$. Initially $[\mathrm I]=0$ can mean zero <reaction product> formation while A already disappears. In particular, $-d[\mathrm A]/dt=d[\mathrm P]/dt$ requires negligible intermediate accumulation here. A single <elementary reaction> with the same net <chemical equation> would instead predict immediate formation at $k[\mathrm A][\mathrm B]$. Measuring both sides can reveal a distinction hidden by the net equation.
This makes <reaction mechanisms> useful for designing observations: follow a candidate <reaction intermediate>, measure individual <reaction products>, or vary one starting <concentration> at a time. The <consecutive first-order reaction mechanism> quantifies an intermediate's rise and fall; the <parallel first-order reaction mechanism> shows how identical <reactant> traces can conceal different <reaction products>. The <acetone iodination rate-law experiment> connects such reasoning to a real laboratory protocol. The existing <reaction coordinate>, <potential energy surface>, and <energy profile (chemistry)> describe configurations and barriers along a route; a <concentration> plotted against time describes a different observable.
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