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Reaction Mechanisms for AP Chemistry

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By — McGraw-Hill Professional
Updated on Feb 2, 2011

Practice problems for these concepts can be found at:

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 reactio

    A + 2B → E + F

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

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.

The rate equation for an elementary step can be determined from the reaction stoichiometry, 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:

    NO2(g) + CO(g) → NO(g) + CO2(g)

It has been determined experimentally that the rate law for this reaction is: Rate = k[NO2]2. This rate law indicates that the reaction does not occur with a simple collision between NO2 and CO. A simple collision of this type would have a rate law of Rate = k[NO2][CO]. The following mechanism has been proposed for this reaction

    NO2(g) + NO2(g) → NO3(g) + NO(g)
    NO3(g) + CO(g) → NO2(g) + CO2(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[NO2]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 reactions a 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.

Practice problems for these concepts can be found at:

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