FIGURE 12.19 Aircraft measurements of CIO and O, on (a) August 23 and (b) September 16, 1987, as the aircraft flew south (adapted from Anderson et al., 1991).

FIGURE 12.20 Vertical profile of CIO on September 19-20, 1992, at McMurdo Station, Antarctica (adapted from Emmons et al., 1995).

FIGURE 12.20 Vertical profile of CIO on September 19-20, 1992, at McMurdo Station, Antarctica (adapted from Emmons et al., 1995).

rates during development of the ozone hole are ~3-4% per day when light is available all day for photolysis (Rosen et al., 1993).

Figure 12.20 shows measurements of the vertical profile of CIO at McMurdo Station, Antarctica, on September 19-20, 1992, as the ozone hole was developing (Emmons et al., 1995). As expected based on the foregoing chemistry, the concentration peaks in the 15-to 20-km range, in the same region as the greatest depletion of 03 is observed. (The peak at higher altitudes is that normally observed globally in the stratosphere due to gas-phase chemistry.)

In short, the overall features of the chemistry involved with the massive destruction of ozone and formation of the ozone hole are now reasonably well understood and include as a key component heterogeneous reactions on the surfaces of polar stratospheric clouds and aerosols. However, there remain a number of questions relating to the details of the chemistry, including the microphysics of dehydration and denitri-fication, the kinetics and photochemistry of some of the C10x and BrOx species, and the nature of PSCs under various conditions. PSCs and aerosols, and their role in halogen and NOx chemistry, are discussed in more detail in the following section.

5. Polar Stratospheric Clouds (PSCs) and Aerosols a. Nature of Aerosols and PSCs

Sulfate aerosol particles with diameters typically in the 0.1- to 0.3-/Lim range are well known to be formed in the stratosphere from a number of sources, forming what is known as the Junge layer. Carbonyl sulfide, COS, is produced in the troposphere by both natural and anthropogenic processes (Chin and Davis, 1993, 1995). It reacts with OH, but the reaction is quite slow, with a room temperature rate constant of 1.9 x 10"15 cm3 molecule-1 s_l (DeMore et al., 1997); this corresponds to a calculated tropospheric lifetime with respect to this one reaction of approximately 17 years at an OH concentration of 1 x 10h radicals cm"3. Based on measured atmospheric concentrations of COS and estimated source strengths, Chin and Davis (1995) estimate a global atmospheric lifetime of about 4 years. As a result of this lifetime, significant quantities of COS reach the stratosphere, where it is ultimately oxidized to sulfuric acid (Crutzen, 1976; Kourtidis et al., 1995; Zhao et al., 1995). In the absence of volcanic injections of S02, this is the major source of stratospheric sulfate aerosols (SSA) (Crutzen, 1976).

However, the eruption of large volcanoes also injects large quantities of S02 into the stratosphere, increasing the concentration of SSA significantly. For example, typical number concentrations of SSA are about 1-10 particles cm-3 under nonvolcanically perturbed conditions; the number density increases by 1-2 orders of magnitude after major volcanic eruptions (e.g., see Russell et al., 1996).

These aerosols play a major role in stratospheric chemistry by directly providing surfaces for heterogeneous chemistry (discussed in more detail later) as well as serving as nuclei for polar stratospheric cloud formation. Figure 12.21 schematically shows the processes believed to be involved in PSC formation. The thermodynamic stability of the various possible forms of PSCs at stratospherically relevant temperatures and the transitions between them are discussed in detail by Koop et al. (1997a).

The concentration of sulfuric acid in SSA is typically 50-80 wt% under mid- and low-latitude stratosphere conditions. However, as the temperature drops, these particles take up increasing amounts of water, which dilutes the particles to as low as 30 wt% H2S04. Gaseous nitric acid is also absorbed by these solutions, forming ternary H2S04-H20-HN03 solutions with as much as 30 wt% in each acid.

As more and more HN03 and H20 are taken up into solution from the gas phase, the relative amount of H2S04 diminishes until the particle is primarily an HN03-H20 mixture. Continued reduction in the temperature results in nitric acid and sulfuric acid hydrates freezing out. Based on laboratory studies, it has been proposed that nitric acid trihydrate (NAT) freezes out of solution first (Hanson and Mauersberger, 1988a, 1988b; Molina et al., 1993; Iraci et al., 1994, 1995; Beyer et al., 1994), although some studies suggest that much lower temperatures, ~ 170 K, would be required for this to be sufficiently fast in the stratosphere (Bertram and Sloan, 1998b). It has also been proposed that other hydrates such as nitric acid dihydrate (NAD), which nucleates rapidly at stratospheric temperatures from 2:1 H20:HN03 solutions (Tisdale et al., 1997), are formed. Nitric acid pentahydrate (NAP), ternary hydrates such as H2S04 • HN03 • 5H20, or higher hydrates in the form of a water-rich metastable solid phase (vide infra) may also be formed as intermediates prior to the formation of the more stable NAT (Tolbert and Middlebrook, 1990; Marti and Mauersberger, 1993; 1994; Worsnop et al., 1993; Fox et al., 1995; Tabazadeh and Toon, 1996).

Sulfuric acid tetrahydrate (SAT) also ultimately freezes out of these ternary solutions (Molina et al., 1993; Iraci et al., 1995). At higher temperatures found at higher altitudes in the middle and low latitudes, sulfuric acid monohydrate (SAM) may also be stable (Zhang et al., 1995).

In a number of laboratory studies (Molina et al., 1993; Iraci et al., 1994, 1995; Beyer et al., 1994; Kolb et al., 1995), these crystallizations had been observed to form the solid nitric acid and sulfuric acid hydrates a few degrees above the ice frost point, defined as the temperature at which the air is saturated with respect to the formation of a plane surface of ice. [See Marti and Mauersberger (1993) for the vapor pressure of ice at stratospherically relevant temperatures.] However, it appears that this does not occur to a significant extent in the atmosphere. Thus, Carslaw et al. (1994) and Koop et al. (1995, 1997b) report studies showing that ternary HN03-H2S04-H20 solutions do not freeze above the frost point; they suggest that these solutions remain liquid in the stratosphere until the temperature falls below the frost point, where the ice crystals formed act as nuclei for the crystallization of nitric and sulfuric acid hydrates. Similarly, Anthony et al. (1997) followed aerosols composed of solutions of sulfuric and nitric acids and water as a function of time in a low-temperature chamber using FTIR and observed that they remained as supercooled liquids for the duration of the experiments, up to 3 h.

The nitric acid concentration may be a major determinant of the extent of supercooling that occurs for these ternary mixtures (Molina et al., 1993; Song, 1994). For example, Molina et al. (1993) observed that HN03 did not affect the supercooling of H2S04-H20 mixtures at temperatures above 196 K, but below this temperature, the presence of HN03 rapidly promoted freezing. In addition, the availability of seed crystals to promote crystallization appears to be a critical issue. As discussed in detail by MacKenzie et al. (1995), a variety of potential seed crystals and/or surfaces that

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