Reaction Mechanisms

A given set of reactants sometimes have choices as to how they will react and what products will be formed.  Reaction mechanism studies look at how the rearrangement of atoms in a reaction actually occurs.  A reaction mechanism is a series of elementary reactions that is proposed to account for the rate law(kinetics) of a particular reaction.  Pictured below is the energy diagram for a two step reaction.  An  elementary reaction is an individual reaction step in a reaction mechanism. This figure shows two transition states and a total of 2 elementary reactions with the first step being the rate limiting step..  This is an exothermic reaction since the products have lower energy than the reactants.

An example of an elementary reaction is the reaction between the hydrogen radical and chlorine gas to form hydrochloric acid and a chlorine radical :

                                                                         H  + Cl₂ HCl  + Cl

Elementary reactions can be unimolecular -- a single reactant changing into products -- or bimolecular -- two molecules or free atoms forming a new product.  The unimolecular decomposition of ozone is one of the most important reactions in the complex series of reactions that occurs in the upper atomosphere:

                                                                O₃ (g) + UV photon   O₂ (g) + O (g)

Similarly,  an ozone reaction in which elemental oxygen is formed is an  example of a bimolecular reaction.  In this case, a single oxygen atom collides with ozone with sufficient energy to produce two atoms of oxygen gas:

                                                                    O (g)  +  O₃ (g)   2 O₂( g)

In the case of these two reactions, O(g) is the intermediate state produced during one step and consumed during the second step in the reaction mechanism. The rate determining step is the slowest step and controls the rate of the overall reaction.  The overall equation will be the same as the rate determing step equation. Often, in describing reaction mechanisms, it is helpful to discuss the various pathways or "options" that  apply to a set of reactants.  There are numerous strategies for showing the individual steps of a particular reaction mechanism or pathway.  Depending on the environment, the reaction will proceed in a way that favors a low energy state -- in other words, the reactants want to reach equilibrium in the easiest manner possible. Often, the  the intermediate products -- those products formed along the reaction pathway -- are as important as the final products.  The steps in a proposed mechanism must add up to the overall reaction, be physically reasonable and conform to the overall rate law. If a fast step precedes the slow step, the fast step reaches equilibrium, and the concentration of intermediates in the rate law of the slow step must be expressed in terms of reactants.

EXAMPLE PROBLEM:  The following mechanism has been proposed for the formation of N2O5(g) from NO2(g) and O3(g) in the gas phase within clouds: NO2(g) + O3(g)  NO3(g) + O2(g) NO3(g) + NO2(g) + M  N2O5(g) + M Using this information:  What is the overall chemical reaction?  What is the intermediate?  If the experimentally determined rate law for the overall chemical reaction is:   -d[NO2(g)] / dt = k[NO2(g)] [O3(g)] What can be concluded about the relative speeds of each of the two reactions?  The notation indicates the change in the concentration of NO2 as happens over change in time (dt) being equal to the concentrations of the products multiplied times the rate constant.  The negative sign suggests that NO2 is being consumed in the reaction.

SOLUTION:  What is the overall chemical reaction? If you add the two reactions you will obtain : NO2 + O3 + NO3 + NO2 + M  NO3 + O2 + N2O5 + M reduces to: 2 NO2 + O3 N2O5 + O2 to yield the overall reaction.  What is the intermediate? Since NO3 is formed in the first reaction and then consumed in the second reaction, NO3 is the intermediate product of the reaction.  What is the rate law for each step? The first reaction is a bimolecular reaction, and its rate law is second order: -d[NO2] / dt = k1[NO2] [O3] For the second reaction: -d[NO3] / dt = k2[NO3][NO2] and the rate law for the overall reaction is: -d[NO2] / dt = k[NO2] [O3] which is the same as for the first reaction. What can be concluded about the relative speeds of each of the two reactions?   Since the rate law for the overall reaction is the same as the rate law for the first reaction, the first reaction controls the speed of the overall reaction.  This is otherwise known as the "rate-controlling" step or the "rate-determining" step.