In a stepwise reaction, not all bonds are broken and formed at the same time. Hence, intermediates appear in the reaction pathway going from the reactants to the products. A stepwise reaction distinguishes itself from an elementary reaction in which the transformation is assumed to occur in a single step and to pass through a single transition state.
Many other terminologies are used for stepwise reactions: overall reaction, global reaction, apparent reaction, operational reaction, complex reaction, composite reaction, multiple step reaction, multistep reaction, etc.
In contrast to elementary reactions which follow the law of mass action, the rate law of stepwise reactions is obtained by combining the rate laws of the multiple elementary steps, and can become rather complex. Moreover, when speaking about catalytic reactions, the diffusion may also limit the reaction. In general, however, there is one very slow step, which is the rate-determining step, i.e. the reaction doesn't proceed any faster than the rate-determining step proceeds.
- Deprotonation next to (α to) the carbonyl: HC–C=O → C=C–O–
- Attack of enolate: Rδ+ + C=C–O– → R–C–C=O
Reaction intermediates may be trapped in a trapping reaction. This proves the stepwise nature of the reaction and the structure of the intermediate. For example, superacids were used to prove the existence of carbocations.
- Chemical reaction
- Elementary reaction
- Rate equation
- Rate-determining step
- Steady state approximation
- Chemical kinetics
- Lindemann mechanism
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