Master AP Chemistry Topic 5.4 with this free practice quiz. Test your skills on unimolecular, bimolecular, termolecular reactions, and direct rate laws to prepare for your AP exam.
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Elementary Steps and Molecularity
An elementary step is a single molecular event that happens exactly as its equation is written, with no hidden sub-steps in between. A reaction mechanism is built from elementary steps, and the steps added together give the overall reaction.
Molecularity counts the reactant particles that take part in one step. One reactant particle breaking apart or rearranging, such as O₃ → O₂ + O, is unimolecular. Two particles colliding, such as NO + O₃ → NO₂ + O₂, is bimolecular. Three particles colliding at the same instant is termolecular.
Termolecular steps are extremely rare, and steps needing four or more particles effectively never happen: one collision must deliver enough energy and the correct orientation for every particle at once. This is why 2H₂ + O₂ → 2H₂O cannot be a single elementary step even though it balances.
Molecularity counts reactant particles only. Products play no part, and for any elementary step the molecularity equals the overall order of that step’s rate law.
Writing the Rate Law of a Step
For an elementary step, and only for an elementary step, the rate law is read straight from the equation: each reactant’s order equals its coefficient in the step.
The step A + 2B → products has rate = k[A][B]², first order in A, second order in B, third order overall. The step 2NO → N₂O₂ has rate = k[NO]². The step NO₂ + CO → NO + CO₂ has rate = k[NO₂][CO], because both coefficients are 1.
This works because the step is the actual molecular event: the equation states how many particles must collide, and the rate depends on how often exactly that collision occurs. Doubling [B] in the first example makes an A meeting two B particles four times as likely, which is the physical origin of the squared term.
The contrast with overall reactions is absolute. An overall rate law must be measured experimentally and cannot be written from its balanced equation, because the overall equation hides the steps.
Intermediates and Catalysts
An intermediate is a species formed in one elementary step and consumed in a later step. It is a real molecule, but because it is both produced and used up inside the mechanism, it cancels when the steps are added and never appears in the overall equation.
A catalyst inside a mechanism is consumed in an early step and regenerated in a later step. It also cancels from the overall equation, but the timing is reversed: a catalyst is present at the start and restored by the end, while an intermediate does not exist at the start and is gone by the end. Appears first as a reactant and later regenerated: catalyst. Appears first as a product and later consumed: intermediate.
Worked Examples
Example 1: rate law and rate for a bimolecular step
The elementary step 2NO → N₂O₂ has rate constant k = 3.0 M⁻¹ s⁻¹. Find the rate when [NO] = 0.10 M, and the molecularity.
Two NO particles collide, so the step is bimolecular. Its rate law is rate = k[NO]², so rate = 3.0 × (0.10)² = 3.0 × 0.010 = 0.030 M/s. The units of k, M⁻¹ s⁻¹, confirm a second-order step.
Example 2: reading a two-step mechanism
Step 1: NO₂ + NO₂ → NO₃ + NO (slow)
Step 2: NO₃ + CO → NO₂ + CO₂ (fast)
NO₃ is produced in step 1 and consumed in step 2, so NO₃ is the intermediate. Adding the steps and cancelling NO₃ gives the overall reaction NO₂ + CO → NO + CO₂. Step 1 is slow and elementary, so its rate law is written directly: rate = k[NO₂]². With k = 0.50 M⁻¹ s⁻¹ and [NO₂] = 0.20 M, rate = 0.50 × 0.040 = 0.020 M/s.
Traps Specific to This Topic
Applying the coefficient rule to an overall equation is the central error of this topic. If the equation describes a whole reaction rather than one step, its coefficients are not orders.
Molecularity is counted on the reactant side only. The step O₃ + Cl → O₂ + ClO is bimolecular because two reactant particles collide, even though it also makes two product particles.
Catalysts and intermediates are told apart by when they appear, not by whether they cancel. Both cancel from the overall equation. A species consumed first and regenerated is the catalyst; a species produced first and consumed is the intermediate. Swapping the labels is a standard wrong answer.
Never include a product in a step’s rate law. The rate of the step NO₂ + CO → NO + CO₂ depends on [NO₂] and [CO], not on [NO] or [CO₂].
Reject implausible steps on sight. A proposed step with four reactant particles is not elementary, however neatly it matches a measured rate law.