Reaction Mechanisms

The mechanism of a reaction is the actual series of steps through which a chemical reaction occurs.
Consider the reaction below:

Overall Reaction:   A 2     +   2B  →   2AB

This equation seems to imply some sort of encounter between two molecules or atoms of B and one
molecule of A 2 to form two molecules of AB. But suppose instead that the reaction actually takes
place in two steps.

Step    1: A 2  +   B   →   A 2 B (Slow)

Step    2: A 2 B    +   B   →   2AB (Fast)

Note that these two steps add up to the overall (net) reaction. A 2 B, which does not appear in the
overall reaction because it is neither a reactant nor a product, is called an intermediate. Reaction
intermediates, unlike activated complexes, are real molecules that exist at least for a while.
Nonetheless, they may still be difficult to detect.

The slowest step in a proposed mechanism is called the rate-determining step (or the rate-limiting
step), so called because as the bottleneck in the progression of the reaction, it determines the rate
by imposing an upper limit on how fast it goes. In the reaction mechanism given above, for example,
the first step, which has been described as slow, is the rate-determining step; the overall reaction
simply cannot occur any faster than this step, in the same way that in a family outing, no matter
how fast the other members are, everyone will still have to wait for the slowest one to be ready
before the family can set off.

In the discussion on the potential energy diagram above, we have limited ourselves to considering
reactions occurring in a single step. For one that involves several steps, the graph will go through a
series of “hilltops and valleys”: Each step will involve a transition state. The low points between