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l 6 Heterogeneous Chemical Processes in Ozone Depletion MARIO J. MOLINA Jet Prop vision Laboratory National Aeronautics and Space Administration This discussion of heterogeneous chemistry focuses first on the reaction between hydrogen chloride (MCI) and chlorine nitrate CONGO. By the term Heterogeneous chemistry" is meant a reac- tion that does not occur only in the gas phase but instead requires a condensed phase in order to proceed. As the previous speakers have indicated, free chlorine atoms are believed to be the principal agent for destroying ozone. Hydrogen chloride and chlorine nitrate species act as sinks, or reservoirs, of chlorine in the atmosphere. These species are not directly reactive with ozone. Hydrogen chloride is a well-known species with well-known properties. Chlorine nitrate is more esoteric, requiring a description of some of its properties. Chlorine nitrate has two peculiar properties in the context of atmospheric chemistry. One is that in contrast to most other species containing chlorine and composed of many atoms, it absorbs light relatively inefficiently; hence it is an efficient sink for storing chIo- rine. Related species with fewer oxygen atoms photolyze much more readily and do not serve as efficient reservoirs of chlorine. The other interesting property of chlorine nitrate is that it is a very difficult species to synthesize in the laboratory. On the surfaces of reac- tion vessels, it is chemically unstable and decomposes rapidly. Thus it came as a surprise to many atmospheric chemists that it was a comparatively stable chemical in the stratosphere. If it reacts with 48

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HETEROGENEOUS OH~ICAL PROCESSES . 49 hydrochloric acid on a suitable surface, then the molecular chlorine gas that is produced is quickly broken down into chlorine atoms by the action of absorbed light. In the context of antarctic chemistry, when this reaction occurs on ice surfaces, the other product of the reaction, nitric acid, re- mains bound to the ice. The atomic chlorine, on the other hand, reacts with ozone to produce chlorine monoxide and molecular oxy- gen, destroying two ozone molecules in the process. The chlorine monoxide combines with nitrogen dioxide to reform chlorine nitrate, thereby partially replenishing the chlorine reservoir and removing any remaining nitrogen dioxide from the air. The net effect of the cycle is to release chlorine from the hydrogen chloride (normally very stable), to scavenge nitrogen from the atmosphere, and to generate chlorine monoxide while simultaneously converting ozone to molecu- lar oxygen. Before the discovery of the antarctic ozone hole, the reaction between hydrogen chloride and chlorine nitrate was considered to be improbable for several reasons. The principal surfaces believed to be effective in the stratosphere in promoting the reaction were those of sulfuric acid, but they appeared to be only weakly favorable for making the reaction proceed. In adclition, a simultaneous collmion of two gaseous molecules on a surface appeared to be necessary in order for the reaction to occur. Such a simultaneous collmion has a very low probability of occurring. In 1986, we began to study reactions that could be taking place in the polar night stratosphere over Antarctica. We and others surmised that stratospheric ice clouds relight be important in causing reactions, since such stratospheric clouds are found almost exclusively in the Antarctic in winter. When it does get cold enough for ice clouds to form, the crystals are not just pure ice. The conclusions of other researchers were that nitric acid would be bound to the solid ice crystal, even at temperatures a few degrees above freezing, through the formation of various hydrates of nitric acid, but that hydrochloric acid would not condense and hence not be bound. S. Pickering (1893) published a study on the formation of pure Ice and of hydrogen chloride hydrate crystals at low temperatures by freezing aqueous solutions. The result of his study was that the concentration by weight of hydrogen chloride compared to water must be greater than about 24 percent for any of it to be incorporated into the frozen crystalline form. This would imply that hydrogen chloride at usual atmospheric concentrations would have no affinity for ice.

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l 50 C) o - MARIO J. MOLINA o -20 -40 -60 -80 -100 ~1 1 1 ~ _ \\ \ LIQUID ~ -\\ \ err ~ \ /HCQ-3H2O - \\ \' / SOLID) \ / _ (ICE) l ~ / SOLID ~\ / (HYDRATES) _ ~ V 1 1 1 1 1 1 0 20 40 60 % HC] FIGURE 6-1 Equilibrium phase diagram for the hydrogen chloride (HC1) water system as a function of temperature and HC1 concentration. Other studies conducted at temperatures close to freezing came up with the same result. We carried out experiments to check these results. We measured the hydrogen chloride concentration directly in the solid and in the liquid phase at temperatures appropriate to achieve equilibrium. For initially liquid solutions with a concentration of less than 24 percent, the resultant hydrogen chloride concentration in the solid phase was between about one-third and one-fourth of the corresponding liquid- phase value (Figure ~1~. This result, in contradiction to those of the earlier studies, implies that polar stratospheric ice clouds will absorb significant amounts of hydrogen chloride vapor. We concluded that ice is actually very efficient in scavenging gas-phase hydrogen chloride. We therefore expect that most of the hydrogen chloride in the antarctic stratosphere is in the condensed phase and bound to ice crystals when they are present. The nature of

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HETEROGENEOUS CHEMICAL PROCESSES 51 the nitric acid hydrates that are likely to be present in the antarctic ice crystals needs further study, as the chemistry involved appears to be more complicated than was originally thought. Nevertheless, the end result is simply that both hydrogen chloride and nitric acid are efficiently scavenged by stratospheric ice particles. We thus realized that if the hydrogen chloride is already in the ice, a simultaneous collision of two species on a surface is not required to cause the hydrogen chioride-chIorine nitrate reaction to occur. Instead, the collision of a chlorine nitrate gas molecule with an ice particle containing hydrogen chloride, an event with a much higher probability of occurring, might be sufficient. We first carried out experiments to measure diffusion rates for hydrogen chloride in ice at temperatures around 200 K. We used optical absorption techniques to determine the rate of penetration of hydrogen chloride into ice. We were surprised to find that hydrogen chloride moves extremely rapidly within ice. This, we believe, is a re- sult of the relative ~openness" of the ice crystalline lattice, combined with the fact that hydrogen chloride molecules are comparatively small, allowing their rapid movement through the lattice. We con- cluded that hydrogen chloride diffuses almost as quickly in solid ice at 200 K as it does in liquid water, in contrast to the conventional view that diffusion in solids is always much slower. We measured movements of several millimeters on a time scale of just minutes. The next experiment we did was to look at the infrared spectra of ice samples. We compared the spectrum of pure ice with ice that contained a fraction of a percent of hydrogen chloride: it can be used as a diagnostic to determine if there is hydrogen chloride, even in very small amounts, in ice crystals (Figure 6-2~. We did the same experiment with nitric acid. When it is first deposited on ice, the spectrum is very similar to that for condensed nitric acid. After about an hour, the spectrum changes, to show a structure around wave number 700. This spectrum change indicates that the nitric acid is being absorbed by the ice. Finally, the spectrum for ice containing hydrogen chloride and chlorine nitrate is essentially the same as that for ice containing hydrogen chloride and nitric acid (Figure ~3~. This implies that when chlorine nitrate is deposited on an ice crystal that was previously treated with hydrogen chloride, nitric acid is produced that remains in the solid phase within the ice. We also carried out experiments with ice-coated tubes in order to determine the probability of interaction with ice of several gases. For hydrogen chloride, we determined that at least one of every five

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l 52 I1J At: CY lo MARIO J. MOLINA ICE ECU+ ICE I ~ 4500 3500 2500 WAVENUMBERS 1500 500 FIGURE 6-2 Infrared spectra of pure ice compared with that for ice containing approximately 1 percent HCI. (Reprinted, by permission, from Molina et al., 1988. Copyright @1988 by The American Association for the Advancement of Science.) collisions resulted in the hydrogen chloride being scavenged by the ice. For chlorine nitrate, we also found that the "sticking coefficient" is very large if the ice on the surface of the tube is doped with hydrogen chloride. We also measured molecular chlorine as a product of the collisions. The measured reaction rate is sufficient to explain the observations in the antarctic stratosphere. The conclusion of all these experiments is that the chlorine nitrate-hydrogen chloride reaction is very efficient in the presence of ice and produces molecular chlorine, thus converting chlorine from an inactive reservoir form to a form that is readily affected by ultravi- olet (UV) radiation (Figure 6-4~. There are two other reactions that

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HETEROGENEOUS CHEMICAL PROCESSES 53 HNO3 + HCQ + ICES Lo Ad c CO on `_. ,~_ 4:O2 + HCQ + ICE ~4, ~I ~ at, \ i ~ ~1~ / ~ ~ ' 1 1 1 1 1 1 1 4500 4000 3500 3000 2500 2000 1500 1000 500 WAVENUMBERS FIGURE 6-3 Comparison of infrared spectra for ice exposed to small amounts of the following compounds: HCl (approximately 1 percent), HCl together with nitric acid (HNO3), and HCl together with chlorine nitrate (ClONO2~. (Reprinted, by permission, from Molina et al., 1988. Copyright (~1988 by The American Association for the Advancement of Science.) can also result in the "activation" of chlorine. One is the reaction of nitrogen pentoxide (N205) with the hydrogen chloride in the ice to produce nitric acid and nitry} chloride, which is relatively unstable in the presence of UV. This reaction has been studied by the SRI group and also by a group at the Jet Propulsion Laboratory and has been shown to be an important one for chlorine conversion. The second reaction is that of hypochIorous acid (HOCI) with hydrogen chloride to produce molecular chlorine and water. This reaction also proceeds fairly quickly. Thus, the presence of other species besides chlorine nitrate can result in liberation of the chlorine from the hydrogen chloride in the ice. Clearly, the presence of ice crystals in the atmosphere is very fa- vorable for the release of chlorine that can destroy ozone. A question that remains is the relative efficiency of ice that contains nitric acid,

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54 MARIO J. MOLINA HC1 + CloNO2-HNO3 (S) + C12 C12 + he 2C1 2tC1+O3-C10+02] C10 ~ NO2-ClONO2 NET: HC1 + 2O3 + NO2-HNO3(s) + 2O2 ~ C10 [HC1; > ~ [NO2] FIGURE 6-4 Series of reactions that result in the destruction of ozone (03) molecules by chlorine compounds in the presence of polar stratospheric ice clouds. The net result of these reactions is also shown. compared to pure ice, in scavenging hydrogen chloride and promot- ing a reaction with chlorine nitrate. We have done some prelirn~nary studies to answer this question. When we doped the ice with large amounts of nitric acid, the reaction no longer occurred. With pro- gressively smaller amounts of nitric acid, the reaction rate increased. A very similar situation probably obtains with sulfuric acid. The critical point is the amount of water available in the condensed phase for the reaction to occur. Here, we can make a thermodynamic ar- gument: if the temperature is very close to the freezing point of water, even if not to the point where pure ice can be crystallized, the existence of a condensed phase will be sufficient to promote these reactions. Therefore, it seems that any of these acid solutions that is sufficiently close to the frost point of water will promote the reactions in the same way as we have shown. (In answer to a question on the possible role of large volcanic eruptions in producing sulfuric acid-water ice): The extent to which sulfuric acid droplets would be effective in promoting these types of reactions would depend on the available water that the droplets

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HETEROGENEOUS CHEMICAL PROCESSES 55 contain. Concentrated sulfuric acid (95 percent acid) has the small amount of water present so tightly bound that the water is not avail- able for reaction. Fairly dilute solutions may be effective in promot- ing reactions; laboratory experiments will be needed to determine this. A plausible argument is that if sulfuric acid droplets exist at temperatures close to the frost point of water in the surrounding atmosphere, then they will become very dilute. What happens will depend critically on the temperature, and experiments should be conducted with that in mind. (In answer to a question about the effects of methane increase and greenhouse gas cooling on stratospheric water-vapor content): Methane and carbon dioxide are usually considered to be the "good guys" in terms of countering the loss of ozone due to chlorine. But the effect of these gases is the opposite for the kind of heterogeneous chemistry that ~ have discussed. Warming of the surface and lower at- mosphere implies cooling of the stratosphere. Both the temperature decrease and the additional water vapor, formed from methane, that reaches the stratosphere will favor greater formation of stratospheric ice clouds. REFERENCES Molina, M.J., T.-L. Tso, L.T. Molina, and F.C.-Y. Wang. 1988. Antarctic stratospheric chemistry of chlorine nitrate, hydrogen chloride, and ice: release of active chlorine. Science 238:1253-1257. Pickering, S. 1893. Die hydrate der chlorwasserstoffsaure. Ber. Dtsch. Chem. Ges. 26:277-289.