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Molecularity is the number of reactant molecules that come together in a single elementary step: unimolecular (one molecule rearranges or dissociates), bimolecular (two molecules collide), or termolecular (three molecules collide simultaneously, which is rare). For elementary reactions, molecularity directly determines the rate law -- a bimolecular step A + B -> products has rate = k[A][B]. Reaction order, by contrast, is an empirical quantity determined from the overall rate law of the observed reaction, which may involve multiple elementary steps. For complex (multi-step) reactions, the overall order bears no necessary relation to the stoichiometry or to the molecularity of any individual step. The distinction is critical: molecularity is a mechanistic concept that applies only to elementary steps, while order is an experimental observable that applies to the overall reaction.
Examine a multi-step mechanism (e.g., the decomposition of N2O5 or the H2 + Br2 reaction) and derive the overall rate law using the steady-state or pre-equilibrium approximation. Compare the resulting overall order to the molecularity of each individual step to see clearly that they differ.
From collision theory, you know that reactions occur when molecules collide with sufficient energy and proper orientation. Molecularity formalizes this at the level of a single elementary step: it is simply the count of reactant molecules (or atoms, or ions) that participate in that one step. A unimolecular step involves one molecule rearranging or breaking apart on its own (like the isomerization of cyclopropane to propene). A bimolecular step involves two molecules colliding and reacting (like SN2 displacement or an E2 elimination). A termolecular step would require three molecules to collide simultaneously — which is so statistically unlikely that genuine termolecular elementary steps are exceedingly rare.
The crucial distinction is that molecularity applies only to elementary steps — reactions that occur in a single event with no intermediates. For an elementary step, the rate law follows directly from molecularity: a unimolecular step A → products has rate = k[A], a bimolecular step A + B → products has rate = k[A][B], and so on. This is not an empirical observation — it is a logical consequence of the step being elementary. If two molecules must collide for the reaction to happen, the rate must depend on the concentrations of both.
Reaction order, by contrast, is an empirical quantity. It describes how the experimentally measured rate of the overall reaction depends on concentration: if rate = k[A]m[B]n, then the reaction is m-th order in A, n-th order in B, and (m + n)-th order overall. For an elementary reaction, order equals molecularity. But most reactions are not elementary — they proceed through a mechanism of multiple elementary steps, and the overall rate law is determined by the rate-limiting step and the relationships between intermediates. The overall order can be fractional, zero, negative, or any value; it bears no necessary relationship to the stoichiometric coefficients of the balanced equation.
Consider a concrete example: the decomposition of ozone, 2O₃ → 3O₂. The stoichiometry might suggest second order, but the experimentally observed rate law is rate = k[O₃]²[O₂]⁻¹ — the reaction is negative first-order in O₂, something that makes no sense if you try to read order from stoichiometry. The mechanism involves a fast equilibrium (O₃ ⇌ O₂ + O) followed by a slow bimolecular step (O + O₃ → 2O₂). Deriving the rate law from this mechanism, using the pre-equilibrium approximation, yields the observed rate expression. The molecularity of each step is well-defined (unimolecular dissociation, then bimolecular collision), but the overall order reflects the combined kinetics of the entire mechanism. Keeping this distinction clear — molecularity describes mechanism, order describes measurement — is essential for correctly interpreting kinetic data and proposing mechanisms.