Reaction Mechanisms

Most chemical reactions happen as a sequence of multiple reactions, in a step-wise manner known as a reaction mechanism. For example, take the following reaction mechanism:

O₃(g) → O₂(g) + O

O + O₃(g) → 2O₂(g)

We call the individual steps in the mechanism "elementary steps". The sum of these two distinct reactions will give the net chemical reaction that takes place. The net reaction is below(O cancels out):

2O₃(g) → 3O₂(g)

In terms of collision theory, the elementary steps also tell us how the reaction truly proceeds. It's not that 2 ozone molecules yield diatomic oxygen through a collision. Rather, ozone collides to yield oxygen molecules and oxygen atoms. More ozone and oxygen atoms then collide to make diatomic oxygen molecules again.

Elementary reactions with one reactant are called unimolecular, elementary reactions with one reactant are called bimolecular, elementary reactions with three reactants are called termolecular.

In reaction mechanisms, catalysts show up as reactants in the first step and as products in the last step, as they don't get produced or consumed across the span of a total reaction mechanism. When a species is produced and reacted in the reaction mechanism but not in the net reaction, it's known as a reaction intermediate. An example of a reaction intermediate is the oxygen atom O in the reaction mechanism above. It was produced and it reacted during the reaction mechanism but didn't show up in the net reaction.

In a reaction mechanism, the slowest step is the rate-determining step. So, the rate law for a given mechanism uses the concentrations and coefficients of the slowest step. If the slowest step follows a fast reversible reaction, you can substitute any reaction intermediates in the slowest step. You do this by setting the concentration and exponent of that intermediate to whatever else is in the fast equilibrium step. Then, you substitute that into the rate-determining slow step.