Most chemical reactions rarely occur in a single, instantaneous step. Instead, they proceed through a sequence of simpler, microscopic events called the reaction mechanism. Understanding this step-by-step pathway helps predict how quickly a reaction will occur and what factors influence its speed. The simplest of these individual steps are categorized based on how many molecules are physically involved in the transformation.
Defining Molecularity in Chemical Reactions
The distinction between unimolecular and bimolecular reactions centers on molecularity, which refers to the number of reactant species participating in a single, elementary step of a reaction mechanism. Molecularity is always an integer value, representing a count of atoms, ions, or molecules that come together. This count is specific to the slowest step, known as the rate-determining step, which governs the speed of the entire reaction.
A unimolecular elementary reaction involves just one molecule rearranging itself or breaking apart to form products. The single reactant molecule needs only to acquire enough internal energy to overcome the activation barrier and proceed with the transformation. Examples include the thermal decomposition of a molecule or the isomerization of one compound into another.
A bimolecular elementary reaction requires the simultaneous collision of two separate reactant species. These two molecules may be two different compounds, or two molecules of the same compound, but two distinct particles must meet for the reaction to occur. Since a collision is necessary, the rate of a bimolecular step is inherently tied to the concentration of both reactants in the system.
How Reaction Rate Determines Molecularity
While molecularity describes the physical event at the molecular level, its observable effect is measured through the reaction rate. The rate law for a reaction is a mathematical expression that shows how the speed of a reaction depends on the concentration of the reactants. This rate law is the primary indicator of a reaction’s molecularity.
Unimolecular reactions are characterized by a first-order rate law, meaning the reaction rate is directly proportional to the concentration of only one reactant. If the reaction is A \(rightarrow\) Products, the rate law is written as Rate \(= k[text{A}]\), where \(k\) is the rate constant. The reaction rate doubles if the concentration of A is doubled, because twice as many molecules are available to undergo the spontaneous rearrangement.
Bimolecular reactions, conversely, display a second-order rate law, reflecting the requirement for two molecules to collide. If the reaction involves two different reactants, A and B, the rate law is Rate \(= k[text{A}][text{B}]\). Doubling the concentration of either A or B doubles the frequency of collisions between them, thus doubling the reaction rate.
If the bimolecular reaction involves two molecules of the same reactant, the rate law becomes Rate \(= k[text{A}]^2\). In this case, doubling the concentration of A quadruples the reaction rate because the number of possible A-A collisions increases by a factor of four. The exponents in the experimentally determined rate law, which indicate the reaction order, must match the number of molecules involved in the rate-determining step to confirm the molecularity of that step.
Structural Differences in Reaction Pathways
The physical pathway, or mechanism, for each reaction type shows a fundamental structural difference in how bonds change. Bimolecular reactions often proceed through a single, concerted step where bond breaking and bond forming occur simultaneously. This single step passes through a high-energy, transient structure known as a transition state.
The transition state is not a stable, isolatable compound, but rather a fleeting arrangement of atoms that represents the peak of the energy barrier for the reaction. Because two molecules must collide in the correct orientation to reach this single energy peak, the entire transformation happens in one smooth, continuous motion.
Unimolecular reactions frequently involve a multi-step mechanism that features the formation of a distinct, unstable intermediate species. For example, a molecule might first break a bond to form an ion, which is an intermediate, before that intermediate reacts in a second step to yield the final product. This intermediate is a temporary, local minimum on the energy diagram, giving the overall process two or more energy peaks.
This structural difference in the reaction pathway has practical consequences. The formation of charged intermediates in unimolecular reactions can be significantly stabilized by polar solvents, which may accelerate the overall rate. In contrast, bimolecular reactions, which rely on the direct collision of two species, are often less sensitive to solvent polarity but more sensitive to steric hindrance, or the physical bulkiness of the colliding molecules.