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Chemistry 2 • Chemical Kinetics

Elementary Steps & Intermediates — Identify elementary steps, intermediates, and molecularity

Discover how complex reactions break down into simple molecular collisions that reveal the hidden pathway from reactants to products.

Explore the lesson 13 min read
An atomic orbit with a reaction beaker and changing energy trace.
Inside this lesson 10 sections 5 questions 9 figures

SECTION 1

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.

1850s • Early Rate Studies

Ludwig Wilhelmy measured the rate of sucrose inversion in acid, producing one of the first quantitative rate laws and showing that reaction speed depends on concentration.

1884 • Van 't Hoff's Études

Jacobus van 't Hoff published Études de Dynamique Chimique, distinguishing reactions that proceed in one step from those that require multiple sequential stages.

1890s • Concept of Reaction Intermediates

Chemists like Max Bodenstein proposed that unstable species form and then react further during multi-step reactions, giving rise to the modern idea of reaction intermediates.

1935 • Transition-State Theory

Henry Eyring and colleagues developed transition-state theory, providing a mathematical framework for understanding how individual elementary steps proceed through energy barriers.

1960s–Present • Modern Mechanism Studies

Ultrafast spectroscopy and computational chemistry now allow scientists to observe intermediates directly, confirming mechanisms that were once only theoretical proposals.

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.

SECTION 2

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.

1

Elementary Step

A single molecular event that cannot be broken down into simpler reactions. It describes exactly what happens in one collision or bond-breaking event. The rate law for an elementary step can be written directly from its equation.
2

Reaction Intermediate

A species that is produced in one elementary step and consumed in a later step. It does not appear in the overall balanced equation because it cancels out when all elementary steps are added together.
3

Molecularity

The number of reactant molecules (or atoms, ions) that come together in a single elementary step. Molecularity is always a whole number: unimolecular (1), bimolecular (2), or termolecular (3). It applies only to elementary steps, never to overall reactions.
4

Reaction Mechanism

The complete sequence of elementary steps that shows how reactants are converted to products. A valid mechanism must satisfy two requirements: the steps must add up to the overall equation, and the mechanism must be consistent with the experimentally observed rate law.
5

Catalyst vs. Intermediate

A catalyst is consumed in an early step and regenerated in a later step, so it appears in the mechanism but not in the overall equation. Unlike an intermediate, a catalyst is present before the reaction begins.
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KEY TAKEAWAY

Think of a reaction mechanism like driving directions from home to school. The overall equation is just the start and end address. The elementary steps are each individual turn you make. The intermediates are the intersections you pass through but don't stay at—they're important waypoints, but they aren't your starting point or your destination.

SECTION 3

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.

Ozone Decomposition MechanismSTEP 1 — UnimolecularO₃ → O₂ + OOzone breaks apart into molecular oxygen and an oxygen atom (the intermediate).OIntermediateSTEP 2 — BimolecularO₃ + O → 2 O₂A second ozone molecule reacts with the intermediate oxygen atom to form two O₂ molecules.Add steps; O cancelsOVERALL EQUATION2 O₃ → 3 O₂(The intermediate O does not appear here.)LegendIntermediate (formed & consumed)Molecularity: Step 1 = unimolecular (1 reactant); Step 2 = bimolecular (2 reactants)
The diagram traces the two-step ozone decomposition mechanism. Notice how the intermediate oxygen atom (O) is produced in Step 1 and consumed in Step 2. When both steps are summed, the intermediate cancels, yielding the overall equation 2 O3 → 3 O2.

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.

SECTION 4

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.

UNIMOLECULAR RATE LAW

A→productsRate=k[A]A → products Rate = k[A]

One molecule rearranges or decomposes. The rate depends on the concentration of A raised to the first power. Molecularity = 1.

BIMOLECULAR RATE LAW (TWO DIFFERENT SPECIES)

A+B→productsRate=k[A][B]A + B → products Rate = k[A][B]

Two molecules collide. The rate depends on the product of both concentrations. Molecularity = 2.

BIMOLECULAR RATE LAW (TWO OF THE SAME SPECIES)

2A→productsRate=k[A]22A → products Rate = k[A]²

Two identical molecules collide. The rate depends on [A] squared. Molecularity = 2 (count each molecule separately).

TERMOLECULAR RATE LAW (VERY RARE)

A+B+C→productsRate=k[A][B][C]A + B + C → products Rate = k[A][B][C]

Three molecules collide simultaneously. Molecularity = 3. Termolecular steps are exceedingly rare because the probability of three molecules colliding at the same instant is very low.

⚠

Important Distinction

Do not confuse molecularity with reaction order. Molecularity is a theoretical count of how many reactant particles participate in a single elementary step—it is always a positive integer (1, 2, or 3). Reaction order is an experimentally measured exponent in a rate law and can be zero, a fraction, or a negative number. The two are equal only when you are looking at an elementary step.

SECTION 5

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.

MolecularityMeaningExample Elementary StepRate Law
Unimolecular (1)A single molecule decomposes or rearrangesO3 → O2 + ORate = k[O3]
Bimolecular (2)Two molecules collide and reactNO2 + CO → NO + CO2Rate = k[NO2][CO]
Termolecular (3)Three molecules collide simultaneously (very rare)2 NO + O2 → 2 NO2Rate = k[NO]²[O2]

Molecularity categories for elementary steps

How to Identify Intermediates & CatalystsWrite out all elementary stepsList every species on the reactant &product side of each stepAdd all elementary steps to get theoverall equation. Cancel species that appear on both sides.Ask: Was the species present before the reaction started?YESNOCATALYSTConsumed early, regenerated later.Present before & after reaction.INTERMEDIATEProduced in one step,consumed in a later step.Both catalysts and intermediates cancel out of the overall equation.
This flowchart shows a systematic approach to classifying species that cancel from the overall equation. The critical question is whether the species was present before the reaction started. If yes, it is a catalyst; if no, it is an intermediate.

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.

SECTION 6

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.

Identifying Elementary Steps, Intermediates & Molecularity

  1. 1

    Step 1 — Write the proposed elementary steps

    The mechanism is proposed as: Step 1: NO2 + F2 → NO2F + F (slow) Step 2: NO2 + F → NO2F (fast)
  2. 2

    Step 2 — Determine the molecularity of each step

    Count the number of reactant particles in each elementary step. In Step 1, two species collide: NO2 and F2. That makes Step 1 bimolecular. In Step 2, NO2 collides with F, so Step 2 is also bimolecular.
    Step 1: bimolecular; Step 2: bimolecular
  3. 3

    Step 3 — Add the steps to verify the overall equation

    Adding both steps: NO2 + F2 + NO2 + F → NO2F + F + NO2F. The fluorine atom (F) appears on both sides and cancels. The simplified result is 2 NO2 + F2 → 2 NO2F, which matches the given overall equation.
    Overall equation verified ✓
  4. 4

    Step 4 — Identify intermediates

    The fluorine atom (F) was produced as a product in Step 1 and consumed as a reactant in Step 2. It does not appear in the overall equation. Since it was not present before the reaction started, F is a reaction intermediate.
    Intermediate: F (fluorine atom)
  5. 5

    Step 5 — Write rate laws for each elementary step

    Because these are elementary steps, we can write rate laws directly from the equations. Step 1 (slow, rate-determining): Rate1 = k1[NO2][F2]. Step 2 (fast): Rate2 = k2[NO2][F]. Since Step 1 is the slow (rate-determining) step, the overall rate law is approximately Rate = k1[NO2][F2].
    Overall Rate ≈ k₁[NO₂][F₂]

SECTION 7

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.

FeatureElementary StepOverall Reaction
Rate lawCan be written directly from the equation (exponents = coefficients)Must be determined experimentally; exponents ≠ coefficients in general
MolecularityDefined and meaningful (1, 2, or 3)Not defined—molecularity only applies to elementary steps
IntermediatesMay appear as products or reactants in individual stepsDo not appear (they cancel when steps are summed)
ComplexityDescribes a single molecular eventMay involve many hidden steps
Number of particlesRarely more than 2 reactant particles per stepCan have any stoichiometric coefficients

Elementary steps vs. overall reactions

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KEY TAKEAWAY

Think of it this way: the overall equation is like a movie trailer—it gives you the beginning and the end. The elementary steps are the individual scenes that tell you how the plot unfolds. You can't figure out the full story from just the trailer, and you can't write a rate law from just the overall equation. You need the mechanism.

SECTION 8

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 LessonAdvanced Extension
Identifying elementary stepsProposing complete reaction mechanisms from experimental rate laws
Identifying intermediatesApplying the steady-state approximation to derive rate laws that contain no intermediates
MolecularityUsing collision theory and transition-state theory to predict rate constants for each elementary step
Rate-determining stepEnergy profile diagrams showing activation energies for each step and the overall reaction coordinate

How this lesson connects to advanced kinetics topics

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.

SECTION 9

Practice Problems

Question 1: CONCEPTUAL

Explain in your own words why you cannot assign a molecularity to an overall reaction that proceeds through multiple elementary steps.
Show answer
Molecularity describes the number of reactant particles that collide in a single molecular event. An overall reaction is the net result of several such events, so there is no single collision to describe. Saying the overall reaction is 'bimolecular' would be meaningless because the individual steps may each have different molecularities. Only elementary steps—each representing one collision or decomposition—can be assigned a molecularity.

Question 2: BASIC CALCULATION

Consider this elementary step: 2 NO(g) → N2O2(g). (a) What is its molecularity? (b) Write its rate law.
Show answer
(a) Two NO molecules collide, so the molecularity is 2 (bimolecular). (b) Since this is an elementary step, the rate law is written directly from the equation: Rate = k[NO]². The exponent of 2 comes from the stoichiometric coefficient of NO in this elementary step. Remember, this shortcut only works because we are told the step is elementary.

Question 3: INTERMEDIATE

A proposed mechanism for the reaction 2 NO(g) + Br2(g) → 2 NOBr(g) is: Step 1: NO + Br₂ → NOBr₂ (slow) Step 2: NOBr₂ + NO → 2 NOBr (fast) (a) Identify any intermediates. (b) What is the molecularity of each step? (c) Write the rate law predicted by this mechanism.
Show answer
(a) NOBr₂ is the intermediate. It is produced in Step 1 and consumed in Step 2, so it does not appear in the overall equation. (b) Step 1 involves two reactant molecules (NO + Br₂), making it bimolecular. Step 2 also involves two reactant particles (NOBr₂ + NO), so it is also bimolecular. (c) Since Step 1 is the slow (rate-determining) step, the overall rate law is Rate = k₁[NO][Br₂]. This is a second-order rate law that can be verified experimentally.

Question 4: APPLIED

In the catalytic destruction of ozone by chlorine atoms in the atmosphere, the proposed mechanism is: Step 1: Cl + O₃ → ClO + O₂ Step 2: ClO + O → Cl + O₂ (a) What is the overall reaction? (b) Identify any intermediates and any catalysts. (c) Assign the molecularity of each step.
Show answer
(a) Adding both steps: Cl + O₃ + ClO + O → ClO + O₂ + Cl + O₂. Cancel Cl and ClO from both sides to get the overall equation: O₃ + O → 2 O₂. (b) ClO is an intermediate because it is produced in Step 1 and consumed in Step 2, and it was not present before the reaction. Cl is a catalyst because it is consumed in Step 1 and regenerated in Step 2—it was present before the reaction started. (c) Both steps involve two reactant particles, so both are bimolecular.

Question 5: CRITICAL THINKING

A student proposes the following one-step mechanism for the reaction 2 H2(g) + 2 NO(g) → N2(g) + 2 H2O(g): 2 H₂ + 2 NO → N₂ + 2 H₂O (single step) Critique this proposal. Why is it almost certainly incorrect? What rate law would this mechanism predict, and why might that conflict with experiments?
Show answer
If this were a single elementary step, it would be termolecular (or even four-molecule), requiring two H₂ molecules and two NO molecules to collide simultaneously. A four-body collision is essentially impossible under normal conditions—the probability of four molecules all meeting at the same point in space at the same instant is vanishingly small. The predicted rate law would be Rate = k[H₂]²[NO]², which is fourth-order overall. Experiments on this reaction show a different rate law, typically Rate = k[H₂][NO]², supporting a multi-step mechanism instead. This illustrates why most reactions with complex stoichiometry must proceed through multiple elementary steps, each with a molecularity of 1, 2, or at most 3.

SUMMARY

Lesson Summary

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.

Varsity Tutors • Chemistry 2 • Elementary Steps & Intermediates — Identify elementary steps, intermediates, and molecularity