0.15 Reaction rates  (Page 9/10)

The key assumption of the collision model is that the reaction occurs by a single collision. Since thisassumption leads to incorrect predictions of rate laws in some cases, the assumption must be invalid in at least those cases. Itmay well be that reactions require more than a single collision to occur, even in reactions involving just two types of molecules asin . Moreover, if more than two molecules are involved as in , the chance of a single collision involving all of the reactive molecules becomes very small. Weconclude that many reactions, including those in and , must occur as a result of several collisions occurring in sequence, rather than a single collision.The rate of the chemical reaction must be determined by the rates of the individual steps in the reaction.

Each step in a complex reaction is a single collision, often referred to as an elementary process . In single collision process step, our collision model should correctly predict the rate of that step.The sequence of such elementary processes leading to the overall reaction is referred to as the reaction mechanism . Determining the mechanism for a reaction can require gaining substantially more information thansimply the rate data we have considered here. However, we can gain some progress just from the rate law.

Consider for example the reaction in described by the rate law in . Since the rate law involved $\left[N{O}_2\right]^2$ , one step in the reaction mechanism must involve the collision oftwo $N{O}_2$ molecules. Furthermore, this step must determine the rate of the overall reaction. Why would that be? In any multi-step process, ifone step is considerably slower than all of the other steps, the rate of the multi-step process is determined entirely by thatslowest step, because the overall process cannot go any faster than the slowest step. It does not matter how rapidly the rapid stepsoccur. Therefore, the slowest step in a multi-step process is thus called the rate determining or rate limiting step.

This argument suggests that the reaction in proceeds via a slow step in which two $N{O}_2$ molecules collide, followed by at least one other rapid step leading to the products. A possible mechanism is therefore

There are a few important notes about the mechanism. First, one product of the reaction is produced in thefirst step, and the other is produced in the second step. Therefore, the mechanism does lead to the overall reaction,consuming the correct amount of reactant and producing the correct amount of reactant. Second, the first reaction produces a newmolecule, $N{O}_3$ , which is neither a reactant nor a product. The second step thenconsumes that molecule, and $N{O}_3$ therefore does not appear in the overall reaction, . As such, $N{O}_3$ is called a reaction intermediate . Intermediates play important roles in the rates of many reactions.