Why Do Methane and Chlorine Fail to React in the Dark at 25^o?

To reach an understanding of why methane and chlorine do not react in the dark, we must consider the details of how the reaction occurs - that is, the reaction mechanism. The simplest mechanism would be for a chlorine molecule to collide with a methane molecule in such a way as to have chloromethane and hydrogen chloride formed directly as a result of a concerted breaking of the Cl−Cl and C−H bonds and making of the C−Cl and H−Cl bonds (see Figure 4-5). The failure to react indicates that there must be an energy barrier too high for this mechanism to operate. Why should this be so?

Figure 4-5: Four-center collision of chlorine with methane as visualized with ball-and-stick models

First, this mechanism involves a very precisely oriented "four-center" collision between chlorine and methane that would have a low probability of occurrence (i.e., a large decrease in entropy because a precise orientation means high molecular ordering). Second, it requires pushing a chlorine molecule sufficiently deeply into a methane molecule so one of the chlorine atoms comes close enough to the carbon to form a bond and yield chloromethane.

Figure 4-6: Graph of the potential energy of pairs of neon atoms as a function of the internuclear distance. The energy values are per mole of neon atoms.

Generally, to bring nonbonded atoms to near-bonding distances (1.2A to 1.8A) requires a large expenditure of energy, as can be seen in Figure 4-6. Interatomic repulsive forces increase rapidly at short distances, and pushing a chlorine molecule into a methane molecule to attain distances similar to the 1.77-A carbon-chlorine bond distance in chloromethane would require a considerable amount of compression (see Figure 4-7). Valuable information

Figure 4-7: Models showing the degree of atomic compression required to bring a chlorine molecule to within bonding distance of carbon and hydrogen of methane

about interatomic repulsions can be obtained with space-filling models of the CPK type , which have radii scaled to correspond to actual atomic interference radii, that is, the interatomic distance at the point where curves of the type of Figure 4-6 start to rise steeply. With such models, the degree of atomic compression required to bring the nonbonded atoms to within near-bonding distance is more evident than with the ball-and-stick models. It may be noted that four-center reactions of the type postulated in Figure 4-5 are encountered only rarely.

If the concerted four-center mechanism for formation of chloromethane and hydrogen chloride from chlorine and methane is discarded, all the remaining possibilities are stepwise reaction mechanisms. A slow stepwise reaction is dynamically analogous to the flow of sand through a succession of funnels with different stem diameters. The funnel with the smallest stem will be the most important bottleneck and, if its stem diameter is much smaller than the others, it alone will determine the flow rate. Generally, a multistep chemical reaction will have a slow rate-determining step (analogous to the funnel with the small stem) and other relatively fast steps, which may occur either before or after the slow step.

A possible set of steps for the chlorination of methane follows:

Reactions (1) and (2) involve dissociation of chlorine into chlorine atoms and the breaking of a C−H bond of methane to give a methyl radical and a hydrogen atom. The methyl radical, like chlorine and hydrogen atoms, has one election not involved in bonding. Atoms and radicals usually are highly reactive, so formation of chloromethane and hydrogen chloride should proceed readily by Reactions (3) and (4). The crux then will be whether Steps (1) and (2) are reasonable under the reaction conditions.

In the absence of some external stimulus, only collisions due to the usual thermal motions of the molecules can provide the energy needed to break the bonds. At temperatures below 100, it is very rare indeed that the thermal agitation alone can supply sufficient energy to break any significant number of bonds stronger than 30 to 35kcal mol−1.

The Cl−Cl bond energy from Table 4-3 is 58.1kcal, which is much too great to allow bond breaking from thermal agitation at 25o in accord with Reaction (1). For Reaction (2) it is not advisable to use the 98.7kcalC−H bond energy from Table 4-3 because this is one fourth of the energy required to break all four C−H bonds. More specific bond-dissociation energies are given in Table 4-5, and it will be seen that to break one C−H bond of methane requires 104 kcal at 25^o, which again is too much to be gained by thermal agitation. Therefore we can conclude that Reactions (1)-(4) can not be an important mechanism for chlorination of methane at room temperature.

One might ask whether dissociation into ions would provide viable mechanisms for methane chlorination. Part of the answer certainly is: Not in the vapor phase, as the following thermochemical data show:

Ionic dissociation simply does not occur at ordinarily accessible temperatures by collisions between molecules in the vapor state. What is needed for formation of ions is either a highly energetic external stimulus, such as bombardment with fast-moving electrons, or an ionizing solvent that will assist ionization. Both of these processes will be discussed later. The point here is that ionic dissociation is not a viable step for the vapor-phase chlorination of methane.