KAR - Reaction Mechanisms (Lesson)

Reaction Mechanisms

Understanding Reaction Mechanisms

Chemical reactions very often occur in a step-wise fashion, involving two or more distinct reactions taking place in sequence. A balanced equation indicates what is reacting and what is produced, but it reveals no details about how the reaction actually takes place.  A reaction mechanism provides details regarding the precise, step-by-step process by which a reaction occurs.

The decomposition of ozone, for example, appears to follow a mechanism with two steps:

Step #1:  O3 (g)  LaTeX: \longrightarrow\(\longrightarrow\) O2(g) + O
Step #2:  O + O3 (g)  LaTeX: \longrightarrow\(\longrightarrow\)Unknown node type: br2 O2 (g)
Overall Chemical Process:  2 O3 (g) LaTeX: \longrightarrow\(\longrightarrow\)Unknown node type: br3 O2 (g)

As can be seen, the reaction is proposed to occur in a series of two steps known as elementary reactions. These reactions refer to the actual collisions that occur between molecules as opposed to the overall stoichiometry of the reaction. Furthermore, these steps can be described based on the number of molecules participating in each collision. For example, in the mechanism proposed above, Step #1 involves a single reactant particle and can be described as unimolecular. Step #2 on the other hand involves two particles and would be described as a bimolecular event. Termolecular processes involving the simultaneous collision of three particles are exceedingly rare events.

Notice that the monoatomic oxygen atom produced in the first step of this mechanism is consumed in the second step and therefore does not appear as a product in the overall reaction. Species that are produced in one step and consumed in a subsequent step are referred to as intermediates.

Evaluating Reaction Mechanisms

Not only do reaction mechanisms provide insight into the actual collisions that take place during a chemical reaction that eventually leads to the formation of products, but they can also be used to determine rate laws. In order to do so, however, a second piece of information is required.  Refer to the proposed mechanism shown below for the production of ONF:

Overall Reaction:  2 NO (g) + F2 (g)  LaTeX: \longrightarrow⟶  2 ONF (g)

Step #1:  NO + F2  LaTeX: \longrightarrow⟶  ONF + F      Slow

Step #2:  NO + F  LaTeX: \longrightarrow⟶  ONF               Fast

Notice that in this mechanism there is an extra piece of information that is provided. This information allows the determination of the rate-determining step or the slowest step in the overall process. The rate-determining step is essentially the single step in a reaction mechanism that governs the kinetics or rate of the entire process. The rates of the other steps are irrelevant because they are all faster in comparison to this single, slowest step in the process. Because of this, it is possible to write a rate law for the overall chemical reaction from this rate-determining step. For the reaction mechanism shown above it is evident that the first step in the process is the slow step meaning that this is the rate-determining step for the overall process. In addition, because this elementary reaction describes the collisions that occur during this step in the process, it is possible to write a rate law using the coefficients in this reaction. Therefore the rate law for this process becomes:

Rate = k[NO][F2]

It can be seen in this step that the coefficients for both reactants are one; therefore, the rate order with respect to these substances is also one as seen in the rate law. One thing that should be pointed out is that while it is entirely possible to write rate laws using coefficients for reactants in elementary reactions of a reaction mechanism, it is never acceptable to do this when provided an overall chemical equation.  

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