Molecularity

Definition. Molecularity of an elementary reaction is the number of reactant species — molecules, atoms, or ions — that come together at the same moment to bring about that one step's chemical change, exactly as written in its equation. Because it describes a real physical collision, it can only ever be a small positive whole number.

• Unimolecular (molecularity 1). A single molecule transforms on its own — no collision partner needed. Radioactive decay and simple thermal decompositions, such as $\text{N}_2\text{O}_5 \rightarrow \text{N}_2\text{O}_3 + \text{O}_2$ treated as one elementary step, are unimolecular.
• Bimolecular (molecularity 2). Two particles collide and react — by far the most common type, because a two-body collision is geometrically and statistically the easiest event to arrange. Example: $\text{H}_2 + \text{I}_2 \rightarrow 2\text{HI}$ as an elementary step.
• Termolecular (molecularity 3). Three particles must arrive at the same point in space at the same instant, with the right energy and the right mutual orientation. Example: $2\text{NO} + \text{O}_2 \rightarrow 2\text{NO}_2$ as a single elementary step. Such steps are rare, because a genuine three-body simultaneous collision is a much less probable event than a two-body one; most apparent termolecular reactions actually proceed through two successive bimolecular steps instead.

Two rules that make molecularity easy to police. First, molecularity is never zero and never fractional — you cannot have "half a molecule" colliding, so it must be 1, 2, or (rarely) 3. Second, molecularity is defined only for an individual elementary step, never for a complex, multi-step overall reaction; a reaction that proceeds through several steps has no single molecularity of its own, only a molecularity for each of its steps.

Where order fits in. For a genuinely elementary (single-step) reaction, order and molecularity are numerically equal, because the rate law of a true one-step process follows directly from the Law of Mass Action — each reactant's concentration appears in the rate law with an exponent equal to its coefficient in that step. For a complex, multi-step overall reaction, the situation splits: the observed order is fixed by the molecularity of the slowest, rate-determining step, but the overall reaction's experimental order can be fractional, zero, or otherwise disconnected from any coefficient in the balanced equation, while true molecularity — being a physical collision count — never can be.

where: "order" is the sum of the concentration exponents in the experimentally determined rate law, and "molecularity" is the number of colliding species in that one step, as written. Equation (1) holds only when the step in question is truly elementary — it is not a general law connecting order and molecularity for arbitrary reactions.

Molecularity has no formula to derive because it is not a calculated quantity — it is a direct physical description of a single elementary step, read straight off that step's own balanced equation. The reasoning that makes it trustworthy is purely mechanistic, not algebraic, and it comes down to collision probability. A collision between two particles is a common event: molecules in a gas or solution are constantly bumping into each other, so a bimolecular step just needs one such collision to happen with enough energy (matching the activation energy) and the right orientation. A unimolecular step needs even less — a single excited or unstable molecule simply falls apart or rearranges on its own, no partner required. A termolecular step, by contrast, demands that three separate particles converge on the same point in space at the same instant, all three carrying adequate energy and correct mutual orientation simultaneously. The probability of a three-body coincidence like that is dramatically lower than a two-body one, which is exactly why termolecular elementary steps are rare in practice, and why no elementary step of molecularity 4 or higher has ever been established — the odds of four independent particles meeting at once are essentially negligible. This is also why molecularity cannot be assigned to a multi-step overall reaction: the overall balanced equation is only a bookkeeping summary of the net change, built by adding up several elementary steps, and the coefficients in that overall equation need not match the number of particles actually colliding in any real step of the mechanism. Molecularity, being about an actual collision event, only makes physical sense once you have identified the true, single-step mechanism underlying the observed chemistry.

Question. The decomposition $2\text{H}_2\text{O}_2 \rightarrow 2\text{H}_2\text{O} + \text{O}_2$ is experimentally found to be first order overall, with rate law $\text{Rate} = k[\text{H}_2\text{O}_2]$. The balanced equation shows a coefficient of 2 for $\text{H}_2\text{O}_2$. Explain why reading "molecularity = 2" directly from that coefficient would be wrong, and state what can correctly be said about molecularity here.

Step 1. Note the temptation. Seeing the coefficient 2 in front of $\text{H}_2\text{O}_2$ might suggest a bimolecular step where two $\text{H}_2\text{O}_2$ molecules collide at once.

Step 2. Check that temptation against the rate law. If the reaction really were one elementary step with two colliding $\text{H}_2\text{O}_2$ molecules, equation (1) demands order = molecularity = 2, so the rate law would have to be $\text{Rate} = k[\text{H}_2\text{O}_2]^2$. The experimental rate law is instead first order, $\text{Rate} = k[\text{H}_2\text{O}_2]$ — a direct contradiction of the single-step assumption.

Step 3. Apply the rule that molecularity belongs only to elementary steps, not to an overall balanced equation. Since the experimental order does not match the coefficient-based guess, this reaction is not a single elementary step; it must proceed through a multi-step mechanism, and only its individual steps — not the overall equation — can be assigned a molecularity.

Step 4. State the correct conclusion. The overall order of this reaction is 1 (an experimental fact), while the overall reaction itself has no defined molecularity; only once the actual mechanism is known can each elementary step be assigned its own molecularity, and the rate-determining step would have to be unimolecular in $\text{H}_2\text{O}_2$ to reproduce the observed first-order kinetics.

Competitive stretch

Classify each of these as elementary steps and rank them by likelihood: (a) $\text{Cl}_2 \rightarrow 2\text{Cl}^\bullet$; (b) $\text{Cl}^\bullet + \text{H}_2 \rightarrow \text{HCl} + \text{H}^\bullet$; (c) $2\text{NO} + \text{Br}_2 \rightarrow 2\text{NOBr}$. Step (a) is unimolecular (one $\text{Cl}_2$ breaking apart), step (b) is bimolecular (one radical plus one $\text{H}_2$), and step (c) is termolecular (two NO plus one $\text{Br}_2$, three particles total). Ranked by likelihood, unimolecular and bimolecular events (a and b) are far more probable than the termolecular event (c), which needs a simultaneous three-body collision — consistent with why most NO-oxidation-type stoichiometries actually proceed through two successive bimolecular steps rather than one true termolecular step.

• Treating order and molecularity as always equal — equation (1) holds only for a genuinely elementary step, not for a multi-step overall reaction.
• Reading molecularity straight off the coefficients of a complex, multi-step balanced equation — molecularity is defined only for the actual elementary steps of the mechanism.
• Believing molecularity can be zero or fractional — it counts real colliding particles, so it is always a small positive whole number (1, 2, or rarely 3).
• Assuming termolecular steps are common just because a balanced equation shows three reactant particles — they are rare because a genuine three-body simultaneous collision is statistically far less likely than a one- or two-body event.

Ask the tutor: if the order-versus-molecularity distinction still feels blurry, ask your AI tutor for a worked side-by-side comparison table across several multi-step reaction examples.