Getting Started
Most chemical reactions we write as a single balanced equation are actually the sum of several simpler, fundamental steps. This chapter zooms in on the atomic scale to examine these individual steps, known as elementary reactions. We will explore the core process of how individual molecular collisions directly dictate the speed, or rate, of these fundamental chemical events, providing a bridge between the particles we can't see and the reaction rates we can measure.
What You Should Be Able to Do
After completing this section, you will be able to:
Define an elementary reaction as a single step in a chemical process.
Determine the molecularity of an elementary reaction from its balanced equation.
Write the rate law for an elementary reaction using the stoichiometric coefficients of the reactants.
Explain why elementary reactions involving simultaneous collisions of three or more particles are exceptionally rare.
Key Concepts & Analysis
The relationship between reactants and reaction rate in an elementary step is a direct story of cause and effect. The number and type of particles that must collide (the input) directly determines the mathematical form of the rate law (the output). We can analyze this using a process-and-causation framework.
Inputs & Preconditions
The primary input for an elementary reaction is the set of reactant particles—atoms, ions, or molecules—that participate in a single, effective collision. The essential precondition is that the reaction equation represents exactly one such event, not the net result of a multi-step process. The number of particles involved in this single step is a critical property called molecularity.
Unimolecular: One particle undergoes a change (e.g., decomposition).
Bimolecular: Two particles collide to react.
Termolecular: Three particles collide simultaneously to react.
For example, in the elementary reaction O₃(g) → O₂(g) + O(g), the input is a single ozone molecule.
Key Steps / Mechanism
The "mechanism" of an elementary reaction is the collision itself. The rate of the reaction is directly proportional to the frequency of effective collisions between the reactant particles.
Collision Frequency: The more reactant particles there are in a given volume (higher concentration), the more frequently they will collide.
Rate-Concentration Link:
If a reaction requires one molecule of reactant A (
A → Products), its rate depends directly on the concentration of A. Doubling [A] doubles the chance of a molecule reacting, so the rate doubles.Rate ∝ [A]¹.If a reaction requires two molecules of A to collide (
2A → Products), its rate depends on the concentration of A squared. Doubling [A] quadruples the collision frequency between A molecules (A₁ with A₂, A₁ with A₃, etc.), so the rate quadruples.Rate ∝ [A]².If a reaction requires a collision between one molecule of A and one of B (
A + B → Products), its rate depends on the concentration of A and the concentration of B. Doubling [A] doubles the A-B collision frequency, as does doubling [B].Rate ∝ [A]¹[B]¹.
This direct link is unique to elementary reactions. The stoichiometric coefficient of a reactant in an elementary step becomes the exponent (or reaction order) for that reactant in the rate law for that step.
Outputs & Effects
The direct output of this process is the rate law expression for the elementary reaction. This expression mathematically connects the reaction rate to reactant concentrations and a temperature-dependent rate constant, k.
For a unimolecular reaction
A → Products, the rate law is: Rate = k[A]For a bimolecular reaction
2A → Products, the rate law is: Rate = k[A]²For a bimolecular reaction
A + B → Products, the rate law is: Rate = k[A][B]For a termolecular reaction
2A + B → Products, the rate law is: Rate = k[A]²[B]
Controls & Limiting Factors
The primary limiting factor for elementary reactions is probability. While two-particle collisions are very common in a gas or solution, the probability of three specific particles colliding at the exact same instant with the correct orientation and sufficient energy is exceedingly low. The probability of a four-particle simultaneous collision is virtually zero.
This has a profound effect:
Bimolecular reactions are very common.
Termolecular reactions are extremely rare and very slow.
Elementary reactions with a molecularity of four or more are not considered to be plausible steps in any reaction mechanism.
This probabilistic limit is why complex reactions must proceed through a sequence of simpler unimolecular and bimolecular steps rather than one large, multi-particle collision.
Key Models & Representations
The relationship between the molecularity of an elementary step and its corresponding rate law can be summarized in a clear matrix. This model is a powerful tool for translating the stoichiometry of a single reaction step into a predictive rate expression.
| Molecularity | Description | General Form | Example | Rate Law |
|---|---|---|---|---|
| Unimolecular | One particle reacts | A → Products | N₂O₅ → NO₂ + NO₃ | Rate = k[N₂O₅] |
| Bimolecular | Two particles collide | 2A → Products | 2NO₂ → N₂O₄ | Rate = k[NO₂]² |
| A + B → Products | NO₂ + O₃ → NO₃ + O₂ | Rate = k[NO₂][O₃] | ||
| Termolecular | Three particles collide | 2A + B → Products | 2NO + O₂ → 2NO₂ | Rate = k[NO]²[O₂] |
Key Terms, Quantities, & Concepts
Elementary Reaction: A single, fundamental step in a reaction pathway that occurs in one distinct molecular collision or event.
Reaction Mechanism: The complete sequence of elementary reactions that, when combined, describe the overall chemical transformation from reactants to products.
Molecularity: The number of reactant particles (atoms, ions, or molecules) that are involved in the collision described by an elementary reaction. It is always a positive integer (1, 2, or rarely, 3).
Unimolecular: An elementary reaction involving a single reactant particle, having a molecularity of one.
Bimolecular: An elementary reaction involving the collision of two reactant particles, having a molecularity of two.
Termolecular: An elementary reaction involving the simultaneous collision of three reactant particles. These are rare due to the low probability of such an event.
Rate Law: An equation that expresses the relationship between the rate of a reaction and the concentrations of the reactants. For an elementary step, it is derived directly from the reaction's stoichiometry.
Reaction Order: The exponent of a reactant's concentration in the rate law. For an elementary reaction, the order with respect to a reactant is its stoichiometric coefficient.
Skill Snapshots
Causation
Cause: A reaction is defined as an elementary step.
Effect: Its rate law can be written directly from its molecularity, with stoichiometric coefficients becoming the reaction orders.
Cause: A bimolecular step involves the collision of two particles (e.g.,
A + B).Effect: The rate is proportional to the concentration of each particle, leading to a rate law of
Rate = k[A][B].Cause: The probability of a simultaneous three-body collision is extremely low compared to a two-body collision.
Effect: Termolecular elementary reactions are very rare and significantly slower than most bimolecular reactions.
Comparison
Overall Reactions vs. Elementary Reactions: The rate law for an overall reaction must be determined experimentally, while the rate law for an elementary reaction can be determined directly from its stoichiometry.
Molecularity vs. Reaction Order:Molecularity describes the number of particles colliding in a single elementary step (e.g., "a bimolecular reaction"), whereas reaction order is the experimentally determined exponent in the rate law of the overall reaction (e.g., "a second-order reaction"). They are only equivalent for an elementary step.
Bimolecular vs. Termolecular:Bimolecular reactions are common and involve two-particle collisions, while termolecular reactions are rare and involve three-particle collisions.
Change, Continuity, and Time
Baseline Condition: Consider the elementary reaction
NO + O₃ → NO₂ + O₂with initial concentrations of [NO] and [O₃].Change 1: If the concentration of NO is doubled while [O₃] is held constant, the initial rate of this elementary step will double.
Change 2: If the concentrations of both NO and O₃ are doubled, the initial rate of this elementary step will quadruple (2 x 2 = 4).
Continuity: The value of the rate constant, k, for this reaction remains unchanged as concentrations change, but it will change if the temperature is altered.
Common Misconceptions & Clarifications
Misconception: You can always use the coefficients from a balanced chemical equation to write the rate law.
Clarification: This is the most critical mistake to avoid. This shortcut is only valid for elementary reactions. For overall, multi-step reactions, the rate law is determined by the slowest step in the mechanism (the rate-determining step) and must be found through experimental data.
Misconception: The terms "bimolecular" and "second order" are interchangeable.
Clarification: A bimolecular reaction is always second order. However, a second-order reaction is not always bimolecular. An overall reaction could be second order (e.g., Rate = k[A]²) but occur through a multi-step mechanism where the slowest step might be unimolecular or bimolecular. Molecularity describes a single step; reaction order can describe a single step or an entire reaction.
Misconception: All steps in a reaction mechanism occur at the same speed.
Clarification: Reaction mechanisms are composed of multiple elementary steps that can have vastly different rates. The overall rate of the reaction is typically governed by the slowest elementary step, much like the slowest car in a single-lane tunnel sets the pace for all the traffic.
One-Paragraph Summary
Elementary reactions represent the fundamental events of chemical change, occurring as single, discrete molecular collisions. Their defining characteristic is the direct link between stoichiometry and kinetics: the rate law for an elementary step can be written directly from its molecularity, with reactant coefficients serving as the reaction orders. This principle arises because the reaction rate is proportional to the collision frequency of the participating particles. Because the probability of simultaneous collisions decreases dramatically with an increasing number of particles, reaction mechanisms are dominated by common unimolecular and bimolecular steps. Understanding this connection is the first step toward deconstructing complex overall reactions into a series of simple, predictable events.