reactions.The reactions d represent a loss of the intermediates that propagate
the chain reaction. They are called termination steps.All steps can be generally
characterized by the change in the reactive intermediates over the course of the
reaction. Initiation steps form reactive intermediates from reactants, propaga-
tion steps react an intermediate but form another (so there is no net change in
the amount of reactive intermediate), and termination steps decrease the num-
ber of reactive intermediates. A reaction does not have to involve free radicals
to be a chain reaction, although free-radical reactions are the most common
examples of chemical chain reactions.
Many polymerization processes proceed via free-radical mechanisms. Also,
the Cl-atom-catalyzed reactions that contribute to stratospheric ozone depletion
are free-radical reactions that have received a lot of attention in recent years.
Example 20.12
In the discussion of rate-determining steps in section 20.7, the chlorination
of methane, CH 4 , was postulated to occur by a free-radical mechanism:
Cl 2 →2Cl (a)
ClCH 4 →HCl CH 3 (b)
CH 3 Cl 2 →CH 3 Cl Cl (c)
ClCH 4 →HCl CH 3 (d)
CH 3 Cl 2 →CH 3 Cl Cl (e)
a.Classify each reaction a–e as initiation or propagation.
b.No termination step is given above. Suggest some possible termination
steps for this reaction.Solution
a.Since the first reaction creates two free radicals where none existed as re-
actant, reaction a is an initiation reaction. All other reactions have a reactive
intermediate (a free radical) as a reactant and as a product, so reactions b
through e are propagation reactions.
b.A termination reaction reduces the number of reactive intermediates. In
this example, a termination reaction would have two radicals combining to
make a stable molecule. There are three possibilities:
ClCl→Cl 2
ClCH 3 →CH 3 Cl
CH 3 CH 3 →CH 3 CH 3
Experimentally, the presence of a small amount of ethane, CH 3 CH 3 , is seen
as support of the chain reaction mechanism for this chemical reaction.Although all reactions proceed with some change in energy, initiation and
termination reactions usually have an obvious energy change. For initiation re-
actions, there is typically some energy input to promote the formation of re-
active intermediates. In other words, many initiation reactions are endother-
mic. Or, the initiation reaction might occur spontaneously from the molecules
from the higher-energy end of the energy distribution (remember from kinetic
20.9 Chain and Oscillating Reactions 715