5 Steps to a 5 AP Chemistry

Reaction Mechanisms

In the introduction to this chapter we discussed how chemical reactions occurred. Recall that

before a reaction can occur there must be a collision between one reactant with the proper

orientation at the reactive site of another reactant that transfers enough energy to provide the

activation energy. However, many reactions do not take place in quite this simple a way.

Many reactions proceed from reactants to products through a sequence of reactions. This

sequence of reactions is called the reaction mechanism. For example, consider the reaction

A +2B → E +F

Most likely, E and F are not formed from the simple collision of an A and two B mol-

ecules. This reaction might follow this reaction sequence:

A +B →C

C +B →D

D →E +F

If you add together the three equations above, you will get the overall equation A +2B →

E +F. C and D are called reaction intermediates, chemical species that are produced and

consumed during the reaction, but that do not appear in the overall reaction.

Each individual reaction in the mechanism is called an elementary step or elementary

reaction.Each reaction step has its own rate of reaction. One of the reaction steps is slower

than the rest and is the rate-determining step. The rate-determining step limits how fast

the overall reaction can occur. Therefore, the rate law of the rate-determining step is the rate

law of the overall reaction.

The rate equation for an elementary step can be determined from the reaction stoi-

chiometry, unlike the overall reaction. The reactant coefficients in the elementary step

become the reaction orders in the rate equation for that elementary step.

Many times a study of the kinetics of a reaction gives clues to the reaction mechanism.

For example, consider the following reaction:

NO 2 (g) +CO(g) →NO(g) +CO 2 (g)

It has been determined experimentally that the rate law for this reaction is: Rate =

k[NO 2 ]^2. This rate law indicates that the reaction does not occur with a simple collision

between NO 2 and CO. A simple collision of this type would have a rate law of Rate =

k[NO 2 ][CO]. The following mechanism has been proposed for this reaction:

NO 2 (g) +NO 2 (g) →NO 3 (g) +NO(g)

NO 3 (g) +CO(g) →NO 2 (g) +CO 2 (g)

Notice that if you add these two steps together, you get the overall reaction. The first

step has been shown to be the slow step in the mechanism, the rate-determining step. If we

write the rate law for this elementary step it is: Rate =k[NO 2 ]^2 , which is identical to the

experimentally determined rate law for the overall reaction.

Also note that both of the steps in the mechanism are bimolecular reactions, reactions

that involve the collision of two chemical species. In unimolecular reactionsa single

chemical species decomposes or rearranges. Both bimolecular and unimolecular reactions

are common, but the collision of three or more chemical species is quite rare. Therefore, in

developing or assessing a mechanism, it is best to consider only unimolecular or bimolecular

elementary steps.

Kinetics  203