Molecularity of a Reaction

Defining molecularity

The number of reacting species (atoms, ions, or molecules) that must collide simultaneously in an elementary reaction, in order to bring about that chemical change, is called the molecularity of the reaction.

Molecularity is a mechanistic idea — it describes what happens in one elementary step, not the reaction as a whole.

The three kinds of elementary steps

• Unimolecular — a single reacting species is involved, with no collision needed. Example: decomposition of ammonium nitrite, $NH_4NO_2 \rightarrow N_2 + 2H_2O$.
• Bimolecular — the step needs the simultaneous collision of two species. Example: dissociation of hydrogen iodide, $2HI \rightarrow H_2 + I_2$.
• Trimolecular (termolecular) — the step needs the simultaneous collision of three species. Example: $2NO + O_2 \rightarrow 2NO_2$.

The probability that three molecules collide at exactly the same instant is very small, so termolecular steps are rare, and reactions that depend on one are slow to proceed.

Complex reactions and the rate-determining step

Whenever a stoichiometric equation would otherwise require more than three molecules to collide at once, the reaction must actually be occurring across more than one step. A striking example is

which looks, from its equation alone, as though it should be tenth order — but is experimentally found to be second order. This mismatch is itself the evidence that the reaction proceeds through several elementary steps rather than one.

When a reaction has several steps, the overall rate is controlled by whichever step is slowest — just as a relay team's overall time is set by its slowest runner. This slowest step is called the rate-determining step.

A worked mechanism: decomposition of hydrogen peroxide

The decomposition of hydrogen peroxide, catalysed by iodide ion in an alkaline medium,

is found experimentally to follow the rate law

— first order in both $H_2O_2$ and $I^-$. The evidence points to a two-step mechanism:

Both steps are bimolecular elementary reactions. The species $IO^-$ is a reaction intermediate — it is produced and then consumed during the course of the mechanism, but it never appears in the overall balanced equation. Step (1) is the slower of the two, so it is the rate-determining step: the rate at which this intermediate is formed sets the rate of the overall reaction.

Order versus molecularity — keeping the two apart …