Elementary Reactions and Molecularity
An elementary reaction is a single molecular event in a reaction mechanism, and its molecularity (unimolecular, bimolecular or termolecular) is the number of particles that take part in it; for an elementary step, the rate law can be written directly from the equation. In brief: the step NO + O3 → NO2 + O2 is bimolecular, so its rate law is rate = k[NO][O3].
Most reactions in the laboratory do not take place in a single step, but as a sequence of simple events called a mechanism. To understand why a reaction has a certain rate law, one must know the elementary steps. This article explains what an elementary reaction is, how molecularity differs from reaction order, how to check a proposed mechanism and gives examples including a case where two steps add up to an overall equation.
What is an elementary reaction?
An elementary reaction is a reaction that takes place in a single step, with no intermediates, as written in the equation. The particles in the equation collide and rearrange directly into the products. The balanced overall equation shows only the net change, but an elementary step shows what actually happens at the molecular level.
For an elementary reaction (and only for one), the exponents in the rate law are equal to the coefficients of the reactants. This follows from the idea that the rate is proportional to the frequency of collisions between the reacting particles, which is proportional to the product of their concentrations.
What is molecularity?
Molecularity is the number of reactant particles (atoms, molecules or ions) that collide in an elementary step. It is always a positive integer and applies only to elementary steps.
- Unimolecular: one particle changes by itself, such as in isomerization or decomposition. Example: O3 → O2 + O. Rate = k[O3].
- Bimolecular: two particles collide. Example: NO + O3 → NO2 + O2. Rate = k[NO][O3]. For two identical particles, as in 2NO2 → NO3 + NO, rate = k[NO2]2.
- Termolecular: three particles collide at the same time. Example: O + O2 + M → O3 + M, where M is a third body that carries away the excess energy. Rate = k[O][O2][M]. These steps are rare, since the simultaneous collision of three particles is very unlikely.
No steps with a molecularity of four or more are known, because the probability of such a collision is negligible.
Types of elementary steps
| Molecularity | Elementary step | Rate law | Order |
|---|---|---|---|
| Unimolecular | A → products | rate = k[A] | 1 |
| Bimolecular | A + A → products | rate = k[A]2 | 2 |
| Bimolecular | A + B → products | rate = k[A][B] | 2 |
| Termolecular | A + B + C → products | rate = k[A][B][C] | 3 |
How does molecularity differ from order?
| Property | Molecularity | Reaction order |
|---|---|---|
| Applies to | Elementary steps only | Any reaction |
| Found by | Theory (the mechanism) | Experiment |
| Possible values | 1, 2, 3 | 0, 1, 2, 3, fractions, negative |
For a complex reaction, no molecularity can be assigned, and the order may differ from the sum of the coefficients. The order of an elementary step equals its molecularity.
What is a reaction mechanism?
A mechanism is a sequence of elementary steps that leads from the reactants to the products. A proposed mechanism must fulfill two conditions:
- The elementary steps must add up to the overall balanced equation, after the cancelation of the intermediates.
- The rate law derived from the mechanism must agree with the experimental rate law.
Agreement with the rate law does not prove a mechanism, but disagreement shows that it is wrong. Species that are formed in one step and used up in a later step are intermediates. They do not appear in the overall equation.
Worked examples
Example 1: classifying steps
- Cl + O3 → ClO + O2: two particles, so it is bimolecular, rate = k[Cl][O3].
- N2O4 → 2NO2: one particle, so it is unimolecular, rate = k[N2O4].
- 2NO2 → N2O4: two particles, so it is bimolecular, rate = k[NO2]2.
Example 2: a two-step mechanism
The overall reaction is 2NO2 + F2 → 2NO2F, and the experimental rate law is rate = k[NO2][F2]. A proposed mechanism has two steps:
- NO2 + F2 → NO2F + F (slow)
- NO2 + F → NO2F (fast)
Adding the steps: 2NO2 + F2 + F → 2NO2F + F. The F atom cancels (it is an intermediate), which gives the overall equation. The slow step is bimolecular, so the rate law is rate = k1[NO2][F2], in agreement with experiment.
Example 3: why the equation does not give the rate law
The overall equation 2NO + 2H2 → N2 + 2H2O has four reactant molecules. A collision of four molecules at once is practically impossible. The experimental rate law, rate = k[NO]2[H2], shows that the reaction occurs in a series of simpler steps, and the slow step involves two NO molecules and one H2 molecule.
How was the idea developed?
The distinction between the overall reaction and the elementary processes was clarified in the early 20th century. Max Bodenstein’s work on the reaction between hydrogen and bromine in 1906–1907 showed that the observed rate law had a complicated form, which was later explained by a chain mechanism with atoms as intermediates. The idea of elementary steps and of molecularity was refined in the 1920s and 1930s, in particular by Cyril Hinshelwood and Nikolay Semenov, who shared the Nobel Prize in Chemistry in 1956 for their studies of the mechanisms of chemical reactions.
Which mistakes are common?
- Reading the rate law from the overall equation: this works only for elementary steps.
- Assigning a molecularity to a complex reaction: molecularity applies only to elementary steps.
- Forgetting to cancel intermediates: the steps must add up to the overall equation.
- Mixing up molecularity and order: they are equal only for an elementary step.
- Assuming that a mechanism is proved by the rate law: several mechanisms may be consistent with the same rate law.