Chemical reactions transform reactants into products. Although a balanced chemical equation summarizes the overall change, most chemical transformations occur through a series of steps. The elementary reaction is the most basic building block of this sequence, representing a single, indivisible molecular event. Understanding these fundamental steps is necessary for fully describing how a reaction proceeds on a molecular level.
Defining Elementary Reactions
An elementary reaction is a chemical process that occurs in a single, concerted step exactly as it is written. This means the reactants directly collide or rearrange to form products without forming any intermediate species that could be isolated or detected. It is a one-step process characterized by a single transition state, the point of highest energy along the reaction pathway. The absence of detectable species formed and then consumed within the step distinguishes the elementary reaction from more complex overall reactions.
Molecularity: Classifying Reaction Steps
Elementary reactions are categorized by their molecularity, defined as the number of reactant species that participate in that single, concerted step. Molecularity is always a positive whole number, reflecting the number of species that must simultaneously come together for the reaction to occur.
A unimolecular reaction involves only a single reactant species that rearranges or breaks apart to form products, such as thermal decomposition. Bimolecular reactions are the most common type, involving the collision and interaction of two distinct species. These reactions are statistically probable because they only require two molecules to meet with the correct orientation and sufficient energy.
Termolecular reactions, involving the simultaneous collision of three separate species, are exceedingly rare. The probability of three molecules meeting simultaneously with the proper energy and orientation is statistically very low. Consequently, virtually all known chemical reactions are composed of a sequence of unimolecular and bimolecular elementary steps.
The Direct Link to Rate Laws
One important concept for an elementary reaction is the direct relationship between its stoichiometry and its rate law. For any elementary step, the rate law can be written directly by inspection using the reactant stoichiometric coefficients as the reaction orders. The rate law is a mathematical expression showing how the rate of the reaction depends on reactant concentration.
For example, consider the elementary reaction \(A + 2B \rightarrow Products\). The rate law for this specific step must be Rate \(= k[A]^1[B]^2\). Here, the exponents (1 and 2) are the reaction orders with respect to A and B, and they match the stoichiometric coefficients of the balanced elementary equation. The term \(k\) is the rate constant, a proportionality constant specific to that reaction step and temperature.
This direct correlation between stoichiometry and the rate law is a defining feature that applies only to elementary reactions. This contrasts sharply with overall, non-elementary reactions, where the reaction orders must be determined experimentally and frequently do not match the coefficients in the balanced equation.
How Elementary Reactions Fit into Complex Mechanisms
Most reactions encountered in chemistry are complex reactions, meaning they occur through a sequence of two or more elementary steps. This sequential process is called the reaction mechanism, which describes the pathway by which reactants are converted into final products. The sum of all the individual elementary steps within the mechanism must equal the overall balanced chemical equation.
Within a complex mechanism, certain chemical species are formed in one elementary step and then immediately consumed in a subsequent step; these are known as reaction intermediates. The relative speeds of the various elementary steps govern the overall rate of the complex reaction. Typically, one elementary step is significantly slower than all the others, and this step is known as the Rate-Determining Step (RDS).
The RDS acts as a kinetic bottleneck, since the overall reaction cannot proceed any faster than this slowest step. Consequently, the rate law for the entire complex reaction is dictated by the rate law of this single, slowest elementary step.