BOX 1-2

The Rates of Free-Radical Reactions

The rate of a given chemical reaction is affected by a number of parameters, most notably the magnitude of the activation energy required before the reaction can occur. Thus reactions with appreciable activation energies are inherently very slow processes and can often be ignored compared to alternative, faster processes for the chemicals involved. In gasphase reactions involving simple free radicals as reactants, the activation energy exceeds that imposed by their endothermicity by only a small amount. Thus we can assume, conversely, that all exothermic free-radical reactions will have only a small activation energy (Figure la). Therefore, exothermic free-radical reactions usually are fast (providing, of course, the reactants exist in reasonable concentrations in the atmosphere). An example of an exothermic radical energy barrier is

reaction with a small

The activation energy here is only 2 kj/mol.

Reactions involving the combining of two free radicals generally are exothermic, since a new bond is formed, so they too proceed quickly with little activation energy, provided that the radical concentrations are high enough that the reactants do in fact collide with each other at a fast rate.

In contrast, endothermic reactions in the atmosphere will be much slower since the activation barrier must of necessity be much larger (see Figure lb). At atmospheric temperatures, few if any collisions between the molecules would have energy sufficient to overcome this

(continued on p. 50)


Products Reactants Extent of reaction


FIGURE I Potential energy profiles for typical atmospheric free-radical reactions, showing (a) exothermic and (b) endothermic patterns.

BOX 1-2 The Rates of Free-Radical Reactions (continued)

large barrier and allow reaction to occur. An example is the endothermic reaction:

water of a hydrogen atom by ground-state atomic oxygen, given that the reaction is endothermic by about 69 kj/mol. On the same diagram, show the energy profile for the reaction of O* with H20 to give the same products, given that O* lies above ground-state atomic oxygen (O) by 190 kj/mol. From these curves, predict why abstraction by O* occurs quickly but that by O is extremely slow in the atmosphere.

Its activation energy must be at least equal to its AH° = +69 kj/mol, and consequently the reaction would be so very slow at stratospheric temperatures that we can ignore it completely.


Draw an energy profile diagram, i.e., one similar to Figure lb, for the abstraction from

Catalytic Destruction of Ozone by Nitric Oxide

The catalytic destruction of ozone occurs even in a "clean" atmosphere (one unpolluted by artificial contaminants) since small amounts of the X catalysts have always been present in the stratosphere. One important' natural version of X—i.e., one of the species responsible for catalytic ozone destruction in a nonpolluted stratosphere—is the free-radical molecule nitric oxide, NO. It is produced when molecules of nitrous oxide, N20, rise from the troposphere to the stratosphere, where they may eventually collide with an excited oxygen atom produced by photochemical decomposition of ozone. Most of these collisions will yield N2 + 02 as products, but a few of them result in the production of nitric oxide:

We can ignore the possibility that NO produced in the troposphere will migrate to the stratosphere; as explained in Chapter 3, the gas is efficiently oxidized to nitric acid, which is then readily washed out of the tropospheric air, before this process can occur.

The NO molecules that are the products of the above reaction catalyti-cally destroy ozone by extracting an oxygen atom from ozone and forming nitrogen dioxide, NOz; i.e., they act as X in Mechanism I:

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