Molecularity

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. In a unimolecular reaction, a single molecule rearranges atoms, forming different molecules. This is illustrated by the equation A -> P, where \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 where [\rm A] is the concentration of species A, t is time, and 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. 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 which corresponds to the second order rate law: . Here, the rate of the reaction is proportional to the rate at which the reactants come together.