In chemistry, molecularity is the number of molecules that come together to react in an elementary (single-step) reaction and is equal to the sum of stoichiometric coefficients of reactants in the elementary reaction with effective collision (sufficient energy) and correct orientation. Depending on how many molecules come together, a reaction can be unimolecular, bimolecular or even trimolecular. The kinetic order of any elementary reaction or reaction step is equal to its molecularity, and the rate equation of an elementary reaction can therefore be determined by inspection, from the molecularity. The kinetic order of a complex (multistep) reaction, however, is not necessarily equal to the number of molecules involved. The concept of molecularity is only useful to describe elementary reactions or steps.
Unimolecular reactions In a unimolecular reaction, a single molecule rearranges atoms, forming different molecules. This is illustrated by the equation
A ⟶ P , {\displaystyle {\ce {A -> P,}}}
where P {\displaystyle {\rm {P}}} refers to chemical product(s). The reaction or reaction step is an isomerization if there is only one product molecule, or a dissociation if there is more than one product molecule. In either case, the rate of the reaction or step is described by the first order rate law
d [ A ] d t = − k r [ A ] , {\displaystyle {\frac {d\left[{\ce {A}}\right]}{dt}}=-k_{r}\left[{\ce {A}}\right],}
where [ A ] {\displaystyle [{\rm {A]}}} is the concentration of species A, t {\displaystyle t} is time, and k r {\displaystyle k_{r}} is the reaction rate constant. As can be deduced from the rate law equation, the number of A molecules that decay is proportional to the number of A molecules available. An example of a unimolecular reaction, is the isomerization of cyclopropane to propene:
Unimolecular reactions can be explained by the Lindemann-Hinshelwood mechanism.
Bimolecular reactions In a bimolecular reaction, two molecules collide and exchange energy, atoms or groups of atoms. This can be described by the equation
A + B ⟶ P {\displaystyle {\ce {A + B -> P}}}
which corresponds to the second order rate law: d [ A ] d t = − k r [ A ] [ B ] {\displaystyle {\frac {d[{\ce {A}}]}{dt}}=-k_{r}{\ce {[A][B]}}} . Here, the rate of the reaction is proportional to the rate at which the reactants come together. An example of a bimolecular reaction is the SN2-type nucleophilic substitution of methyl bromide by hydroxide ion:
CH 3 Br + OH − ⟶ CH 3 OH + Br − {\displaystyle {\ce {CH3Br + OH^- -> CH3OH + Br^-}}}
Termolecular reactions A termolecular (or trimolecular) reaction in solutions or gas mixtures involves three reactants simultaneously colliding, with appropriate orientation and sufficient energy. However the term trimolecular is also used to refer to three body association reactions of the type:
A + B → M C {\displaystyle {\ce {A + B ->[{\ce {M}}] C}}}
Where the M over the arrow denotes that to conserve energy and momentum a second reaction with a third body is required. After the initial bimolecular collision of A and B an energetically excited reaction intermediate is formed, then, it collides with a M body, in a second bimolecular reaction, transferring the excess energy to it. The reaction can be explained as two consecutive reactions:
A + B ⟶ AB ∗ {\displaystyle {\ce {A + B -> AB}}^{*}}
AB ∗ + M ⟶ C + M {\displaystyle {\ce {AB}}^{*}{\ce {+ M -> C + M}}}
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