Historical Context & Motivation
The study of how fast chemical reactions proceed—chemical kinetics—gained rigor only when scientists realized that most overall reactions are composites of simpler, indivisible molecular events. Before this insight, chemists could measure macroscopic rates but struggled to explain why rate laws so often differed from what balanced equations predicted. The concept of elementary reactions resolved this puzzle by identifying the actual molecular collisions and rearrangements that constitute each step in a reaction mechanism.
The central question that drove this development remains the focus of modern kinetics: if a balanced equation does not directly reveal the rate law, what fundamental molecular events actually control the speed of a reaction? Elementary reactions are the answer—each one represents an irreducible molecular event whose rate law can be written by inspection.
Core Principles & Definitions
An elementary reaction (also called an elementary step) is a reaction that occurs in a single collision or molecular rearrangement event—it cannot be broken into simpler steps. Because the event happens exactly as written, the exponents in its rate law equal the stoichiometric coefficients of its reactants. This is the defining feature that distinguishes elementary reactions from overall reactions, whose rate laws must be determined experimentally.
Molecularity
Rate Law from Stoichiometry
Reaction Mechanism
Rate-Determining Step
Intermediates vs. Catalysts
Visual Explanation — Molecularity
The diagram above illustrates the three categories of elementary reactions classified by molecularity. In a unimolecular step, a single energized molecule undergoes isomerization or bond breaking; the rate law is first order. In a bimolecular step, two species must collide with proper orientation and sufficient kinetic energy, yielding a second-order rate law. Termolecular events—requiring a simultaneous three-body collision with the right geometry and energy—are so improbable that nearly all proposed termolecular reactions are better explained as sequences of bimolecular steps. On the AP exam, you should assume that termolecular steps are essentially nonexistent in proposed mechanisms unless specifically stated otherwise.
Mathematical Framework — Rate Laws from Elementary Steps
The single most important mathematical principle for elementary reactions is that the rate law can be written directly from the balanced elementary step. This is not true for overall reactions. The distinction is critical and appears frequently on the AP Chemistry exam.
When a mechanism has multiple elementary steps, the overall rate law is typically governed by the rate-determining step (RDS)—the slowest step in the sequence. The rate law for the overall reaction is the rate law of the RDS, provided that all species in that rate law are actual reactants (not intermediates). If the RDS rate law contains an intermediate, you must use a prior fast-equilibrium step to express the intermediate's concentration in terms of original reactants.
Reaction Mechanisms — Elementary Steps in Action
A reaction mechanism is a proposed sequence of elementary steps that accounts for the overall stoichiometry and the experimentally observed rate law. Consider the decomposition of ozone in the stratosphere, whose overall equation is 2 O₃(g) → 3 O₂(g). Experimentally, the rate law is rate = k[O₃]²[O₂]⁻¹. A one-step bimolecular collision of two O₃ molecules would predict rate = k[O₃]², which does not match the observed inverse dependence on O₂. The accepted mechanism involves two elementary steps with a reactive intermediate—atomic oxygen, O.
This example showcases several testable concepts. First, elementary steps must sum to give the overall balanced equation. Second, intermediates appear in the mechanism but cancel when steps are added—they never appear in the overall equation. Third, the rate law of the RDS initially contains an intermediate (O), which must be algebraically eliminated using the equilibrium expression from a preceding fast step.
Worked Example — Deriving a Rate Law from a Mechanism
Consider the following proposed mechanism for the reaction 2 NO(g) + Br₂(g) → 2 NOBr(g):
- Step 1 (slow): NO + Br₂ → NOBr₂
- Step 2 (fast): NOBr₂ + NO → 2 NOBr
Determine the overall rate law predicted by this mechanism and identify any intermediates.
Elementary vs. Overall Reactions — Key Distinctions
The most common source of error on kinetics questions is confusing elementary reactions with overall reactions. The table below summarizes the critical differences.
| Feature | Elementary Reaction | Overall Reaction |
|---|---|---|
| Definition | Single molecular event; cannot be broken into simpler steps | Net transformation from reactants to products; may consist of multiple steps |
| Rate law source | Written directly from stoichiometric coefficients | Must be determined experimentally or derived from mechanism |
| Molecularity | Defined (uni-, bi-, or termolecular) | Not applicable — molecularity is undefined for overall reactions |
| Intermediates | May produce or consume intermediates | Intermediates do not appear (they cancel) |
| Order vs. coefficients | Always match | Often do not match |
Connection to Advanced Theory — Transition States & Energy Profiles
Each elementary reaction passes through a transition state (also called an activated complex)—a high-energy, unstable configuration at the peak of the potential energy barrier separating reactants from products. Unlike intermediates, transition states cannot be isolated; they exist for approximately one vibrational period (around 10⁻¹³ s). In a multi-step mechanism, each elementary step has its own transition state and activation energy, and the energy profile shows multiple peaks separated by valleys where intermediates reside.
| Concept | Elementary Reaction Level | Advanced / College Extension |
|---|---|---|
| Activation energy (Eₐ) | Each elementary step has its own Eₐ; the RDS has the largest Eₐ. | Transition state theory relates k to Eₐ via k = (k_BT/h)e^(−ΔG‡/RT). |
| Potential energy diagram | Number of peaks = number of elementary steps. | The shape and height of each barrier depend on the reaction coordinate and molecular geometry. |
| Catalysis | A catalyst provides an alternative mechanism with a lower Eₐ for the RDS. | Enzyme catalysis (Michaelis–Menten) models a two-step mechanism with an enzyme–substrate intermediate. |
| Steady-state approximation | At AP level: fast-equilibrium assumption used to eliminate intermediates. | General steady-state: d[intermediate]/dt ≈ 0, applicable even when equilibrium is not established. |
For the AP exam, you should be able to identify the number of elementary steps from a potential energy diagram (count the peaks), locate intermediates (valleys between peaks), and recognize the rate-determining step as the one with the highest activation energy barrier. In more advanced coursework, transition state theory and the steady-state approximation provide powerful tools for deriving rate laws from complex mechanisms without relying solely on the assumption of a single rate-determining step.