Historical Context & Motivation
For much of chemistry's history, scientists wrote balanced equations and assumed they told the whole story of how a reaction happens. A balanced equation, however, only shows you what goes in and what comes out—it says nothing about the pathway molecules take to get there. By the late 1800s, chemists began to realize that most reactions don't happen in a single dramatic collision. Instead, they unfold through a series of simpler steps, much like a recipe that requires multiple stages before the final dish is complete.
The central question that drove this research was deceptively simple: How do reactions actually happen at the molecular level? A balanced equation might show 2 H2 + O2 → 2 H2O, but it's incredibly unlikely that two hydrogen molecules and one oxygen molecule all slam together at once. Breaking this overall reaction into its elementary steps—and identifying the fleeting intermediates formed along the way—is the key to understanding reaction mechanisms.
Core Principles & Definitions
Before diving into specific examples, you need to master three essential vocabulary terms that form the foundation of reaction mechanisms. These terms describe the building blocks of every multi-step reaction.
Elementary Step
Reaction Intermediate
Molecularity
Reaction Mechanism
Catalyst vs. Intermediate
Visualizing a Reaction Mechanism
The diagram below shows a two-step mechanism for the decomposition of ozone (O3). This is one of the most commonly studied mechanisms in introductory kinetics because it clearly illustrates elementary steps, an intermediate, and molecularity. Study how each step feeds into the next and notice how the intermediate oxygen atom (O) is produced in Step 1 and consumed in Step 2.
There are a few critical observations to take away from this diagram. First, each elementary step describes a single molecular event—ozone breaking apart or ozone colliding with an atom. Second, the oxygen atom (O) labeled as the intermediate never shows up in the overall equation. Third, the molecularity of each step is determined by counting the number of reactant particles in that step alone, not in the overall equation.
Mathematical Framework — Rate Laws from Elementary Steps
One of the most powerful features of elementary steps is that you can write the rate law directly from the balanced equation of each step. This is only true for elementary steps—you can never do this for an overall reaction unless it happens to be a single-step mechanism. For an elementary step, the exponents in the rate law equal the stoichiometric coefficients of the reactants.
Classifying Molecularity & Identifying Intermediates
The table below organizes the three categories of molecularity with examples and expected rate-law forms. Following the table, a second diagram walks you through the systematic process of identifying intermediates and catalysts in any proposed mechanism.
| Molecularity | Meaning | Example Elementary Step | Rate Law |
|---|---|---|---|
| Unimolecular (1) | A single molecule decomposes or rearranges | O3 → O2 + O | Rate = k[O3] |
| Bimolecular (2) | Two molecules collide and react | NO2 + CO → NO + CO2 | Rate = k[NO2][CO] |
| Termolecular (3) | Three molecules collide simultaneously (very rare) | 2 NO + O2 → 2 NO2 | Rate = k[NO]²[O2] |
When you encounter a mechanism problem on a test, follow this flowchart systematically. Write out each step, circle every species that does not appear in the overall equation, and then determine whether each circled species is an intermediate or a catalyst based on whether it was originally present.
Worked Example
Consider the following proposed mechanism for the reaction 2 NO2(g) + F2(g) → 2 NO2F(g). Identify the elementary steps, intermediates, molecularity of each step, and verify the mechanism is consistent with the overall equation.
Comparing Key Concepts — Elementary vs. Overall Reactions
A common source of confusion in kinetics is mixing up properties that apply to elementary steps versus properties that apply to overall reactions. The table below lays out the differences side by side so you can avoid the most frequent exam mistakes.
| Feature | Elementary Step | Overall Reaction |
|---|---|---|
| Rate law | Can be written directly from the equation (exponents = coefficients) | Must be determined experimentally; exponents ≠ coefficients in general |
| Molecularity | Defined and meaningful (1, 2, or 3) | Not defined—molecularity only applies to elementary steps |
| Intermediates | May appear as products or reactants in individual steps | Do not appear (they cancel when steps are summed) |
| Complexity | Describes a single molecular event | May involve many hidden steps |
| Number of particles | Rarely more than 2 reactant particles per step | Can have any stoichiometric coefficients |
Connection to Rate-Determining Steps & Advanced Kinetics
Once you can identify elementary steps and intermediates, the next major concept in kinetics is the rate-determining step (RDS). The rate-determining step is the slowest elementary step in a mechanism, and it acts as a bottleneck for the entire reaction. Just like a highway's traffic flow is limited by the narrowest lane, a reaction can only proceed as fast as its slowest step. The overall rate law is determined by the RDS, which is why identifying elementary steps correctly is so critical.
| Concept in This Lesson | Advanced Extension |
|---|---|
| Identifying elementary steps | Proposing complete reaction mechanisms from experimental rate laws |
| Identifying intermediates | Applying the steady-state approximation to derive rate laws that contain no intermediates |
| Molecularity | Using collision theory and transition-state theory to predict rate constants for each elementary step |
| Rate-determining step | Energy profile diagrams showing activation energies for each step and the overall reaction coordinate |
In college-level courses, you will learn to apply the steady-state approximation, which assumes that the concentration of an intermediate remains approximately constant during the reaction. This technique lets you algebraically eliminate intermediates from rate expressions. For now, the essential skill is recognizing elementary steps, correctly assigning molecularity, and identifying which species are intermediates. These abilities are the foundation upon which all advanced mechanism analysis is built.
Practice Problems
Lesson Summary
A reaction mechanism is the step-by-step sequence of elementary steps that shows exactly how reactants become products at the molecular level. Each elementary step describes a single collision or decomposition event, and its molecularity—the number of reactant particles involved—is always 1 (unimolecular), 2 (bimolecular), or rarely 3 (termolecular). A unique and powerful feature of elementary steps is that their rate laws can be written directly from the balanced equation of that step—the exponents equal the stoichiometric coefficients.
When elementary steps are added together, species that are produced in one step and consumed in another cancel out. If a cancelled species was not present before the reaction began, it is a reaction intermediate. If it was present initially, it is a catalyst. Neither intermediates nor catalysts appear in the overall balanced equation. Understanding these concepts is essential for analyzing the rate-determining step and predicting the overall rate law of a reaction—the central skills of chemical kinetics.