Reaction mechanisms: the steps behind the equation
Learn what reaction mechanisms are, how elementary steps and the rate-determining step work, and how a mechanism must predict the observed rate law.
Quick answer
A reaction mechanism is the step-by-step sequence of elementary molecular events — collisions and rearrangements — by which reactants become products. The slowest step, the rate-determining step, controls the overall rate.
Example: For 2NO + O₂ → 2NO₂, a two-step mechanism is: (1) NO + O₂ → NO₃ (slow), (2) NO₃ + NO → 2NO₂ (fast). The steps sum to the overall equation, and the slow first step predicts rate = k[NO][O₂].
What you will learn
- Define elementary steps, intermediates, and the rate-determining step.
- Verify that a proposed mechanism's steps sum to the overall balanced equation.
- Derive the predicted rate law for a simple mechanism with a slow first step.
- Explain how experimental rate laws test proposed mechanisms.
Mechanisms: the hidden choreography
A balanced equation is a summary, not a story. It says 2NO + O₂ becomes 2NO₂, but it does not say whether three molecules collide at once (essentially impossible) or the change happens in stages. A reaction mechanism proposes those stages: a sequence of elementary steps, each a single molecular event like a collision or a rearrangement.
Two rules constrain every proposed mechanism:
- The elementary steps must add up to the overall balanced equation — every intermediate produced in one step must be consumed in another.
- The mechanism must predict the experimentally observed rate law.
A mechanism that fails either test is wrong, no matter how plausible it looks.
Elementary steps and intermediates
Each elementary step has a molecularity: the number of particles colliding in that step. Unimolecular (one particle rearranging), bimolecular (two colliding), and rarely termolecular (three colliding).
Steps also create intermediates — species formed in one step and destroyed in a later one. In the NO oxidation mechanism below, NO₃ is an intermediate: it appears in step 1 and vanishes in step 2, so it never shows in the overall equation. Intermediates are real but fleeting, and detecting them (by spectroscopy or trapping) is strong evidence for a mechanism.
The rate-determining step: a worked mechanism
In a multi-step sequence, the slowest step — the rate-determining step — sets the pace for the whole reaction, like the narrowest section of a highway setting the traffic flow. Consider:
- Step 1 (slow): NO + O₂ → NO₃
- Step 2 (fast): NO₃ + NO → 2NO₂
Adding the steps: NO + O₂ + NO₃ + NO → NO₃ + 2NO₂, which simplifies to 2NO + O₂ → 2NO₂ — the correct overall equation, with NO₃ canceling as an intermediate.
Because step 1 is slow and bimolecular, the predicted rate law is rate = k[NO][O₂] — first order in each. If experiments confirm this rate law, the mechanism survives; if they show something else, it is back to the drawing board.
Why mechanisms matter
Mechanisms are how chemists debug and design reactions. If a catalyst speeds a reaction, the mechanism shows which step it accelerates. If an unwanted side product appears, the mechanism reveals where the pathway branches.
Proposing mechanisms is constrained detective work: every candidate must satisfy atom balance across its steps and reproduce the measured rate law. Techniques like isotopic labeling — swapping an atom for its isotope and tracing where it ends up — then provide independent confirmation.
Variables and units
| Symbol | Name | Unit | Meaning |
|---|---|---|---|
| Molecularity | unitless | Number of reactant particles colliding in a single elementary step (1, 2, or rarely 3) |
Common mistakes
The overall equation shows only the net change — what you start with and what you end with. The mechanism is the hidden choreography of bond-breaking and bond-forming events in between, and it is usually completely different from what the equation suggests.
The slowest step — the rate-determining step — is the bottleneck that sets the pace, like the slowest hiker setting the group's speed. Fast steps before or after it finish quickly and contribute little to the overall time.
An intermediate is created during the reaction and consumed before the end; a catalyst is present at the start and regenerated at the end. Intermediates never appear in the overall equation, while catalysts cancel out of it.
Real-world applications
Catalyst design: mechanisms reveal exactly which step to accelerate.
Drug synthesis: knowing the mechanism lets chemists suppress unwanted side pathways.
Atmospheric chemistry: ozone-depletion mechanisms identified which pollutants to ban.
Polymer manufacturing: mechanism knowledge controls chain length and branching.
Frequently asked questions
How do I get the rate law from a mechanism?
The slow step's reactants (as written in that elementary step) appear in the rate law, raised to the powers of their stoichiometric coefficients in that step. Fast steps before the slow step can contribute too, via pre-equilibrium — which is why deriving rate laws from mechanisms takes care.
Can a species appear in the mechanism but not in the overall equation?
Yes — intermediates are produced and then consumed within the reaction, so they cancel from the overall equation but appear in the middle of the mechanism. Catalysts also cancel from the overall equation but are present from the start and regenerated at the end.
If my mechanism predicts the right rate law, is it proven?
It strongly suggests it, but proof requires more: isotopic labeling, trapping the intermediate, or spectroscopy. Kinetics alone can rule mechanisms out but rarely proves one uniquely — several mechanisms can predict the same rate law.
What is the difference between molecularity and order again?
Molecularity counts particles in one elementary step (always a whole number); order is the measured exponent in the overall rate law (can be fractional or negative). They match only for single-step reactions.