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FIGURE 12.12 Estimated global annual emissions of CFC-11 and CFC-12 (adapted from World Meteorological Organization, 1995).

78 80 82 84 86 88 90 92 Year

FIGURE 12.13 Concentrations of CFC-11 and CFC-12 in air in the 30°N to 90°N region as a function of time. The different curves represent measurements made at various locations (adapted from WMO, 1995).

78 80 82 84 86 88 90 92 Year

FIGURE 12.13 Concentrations of CFC-11 and CFC-12 in air in the 30°N to 90°N region as a function of time. The different curves represent measurements made at various locations (adapted from WMO, 1995).

2. Lifetimes and Atmospheric Fates of CFCs and Halons

In considering the atmospheric fate of CFCs and halons, it is useful to examine the total atmospheric lifetime of a compound X, rx. This is in effect the time required for a pulse emitted into the atmosphere to decay to 1/e of its initial value (see Chapter 5.A.lc). It can be calculated from

LBurden

ELoss Rate '

where EBurdenalm represents the total amount of X in the atmosphere and ELoss Rate the globally integrated loss rate due to all processes, including reactions, uptake into oceans, wet and dry deposition, etc. The relationship between the total atmospheric lifetime rx and the lifetimes with respect to the individual processes that contribute to the removal of the compound X is given by

where the individual terms on the right side of the equation represent the lifetimes with respect to tropo-spheric reactions, stratospheric chemistry, removal by the oceans, deposition, and any other processes that might contribute, respectively. For a compound that reacts only with OH in the troposphere, for example, Tirop = t()ii = 1 AontOH], where kou is the rate constant for the X-OH reaction. Similarly, if there are no known tropospheric sinks, the atmospheric lifetime is the same as the stratospheric lifetime. (Note that this approach addresses removal from the entire atmosphere and does not take into account a number of factors that may be important for individual species.

For example, some may have short lifetimes and/or nonuniform distributions and hence be sensitive to such factors as the location of the emissions, sunlight intensity, season, etc.)

The lifetime of CFCs in the atmosphere can also be estimated using a mass balance approach. Knowing the atmospheric concentrations of CFCs, one can calculate the total amount in the atmosphere. This amount must be the result of a balance between emissions into, and loss from, the atmosphere. If the emission rates are known, the loss rate required to give the observed atmospheric concentrations can be calculated, and from this, a lifetime obtained. Such calculations may be based on either the absolute atmospheric concentrations of CFCs or, alternatively, the observed relative rates of change in the concentrations.

The chlorofluorocarbons (CFCs) have very long lifetimes in the troposphere. This is a consequence of the fact that they do not absorb light of wavelengths above 290 nm and do not react at significant rates with 03, OH, or N03. In addition to the lack of chemical sinks, there do not appear to be substantial physical sinks; thus they are not very soluble in water and hence are not removed rapidly by rainout. While laboratory studies have shown that some of the CFCs decompose on exposure to visible and near-UV present in the troposphere when the compounds are adsorbed on siliceous materials such as sand (Ausloos et al., 1977; Gab et al., 1977, 1978), the lifetimes for CFC-lf and CFC-12 with respect to these processes have been estimated to be ~540 and f800 years, respectively (National Research Council, 1979). Similarly, an observed thermal decomposition when adsorbed on sand appears to be an insignificant loss process under atmospheric conditions.

As a result, CFCs reside in the troposphere for years and are slowly transported up across the tropopause into the stratosphere, primarily in the tropics as discussed earlier. For example, the estimated lifetime of CFC-11 in the global atmosphere, that is, the time to diffuse to the stratosphere and undergo photolysis, is approximately 40-80 years, with that for CFC-12 being about twice as long (WMO, 1995). Tropospheric losses are negligible, so that this is determined by the time to reach the stratosphere and then to dissociate. The stratospheric lifetimes of relevant compounds have been estimated based on their measured concentrations (e.g., Volk et al., 1997; Avallone and Prather, 1997). For example, using tcfc_n =45 + 7 years, Volk et al. (1997) obtained the following lifetime estimates: tN7() = 122 ± 24 years, rCII =93 + 18 years, rCFC.12 = "87 + 17 years, tcfc_i13 = 100 + 32 years, tca = 32 + 6 years, TahCCh = 34 + 7 years, and tiialon.l2ll = 24 + 6 years.

While there are a variety of other chlorinated organ-ics such as methylchloroform (CH3CC13) that are emitted, these have relatively short tropospheric lifetimes because they have an abstractable hydrogen atom (e.g., see WMO, 1995). For example, while the stratospheric lifetime of methylchloroform is estimated to be 34 + 7 years (Volk et al., 1997), its overall atmospheric lifetime is only 5-6 years, primarily due to the removal by OH in the troposphere (tOI[ ~ 6.6 years), with a much smaller contribution from uptake by the ocean (toccan ~ 85 years) (WMO, 1995).

Since the wavelength distribution of solar radiation shifts to shorter wavelengths with increasing altitude (see Chapter 3.C), the CFCs eventually become exposed to wavelengths of light that they can absorb. Figure 12.14 shows the absorption cross sections of some halogenated methanes from 160 to 280 nm (see also Chapter 4.V). The absorptions become very weak beyond ~240 nm in the case of CFC-ff, and 220 nm in the case of CFC-12. Recall from Chapter 3 that both 03 and 02 absorb radiation in the ultraviolet. Figure 12.15 shows these absorption cross sections for 02 and 03 from f20 to 360 nm; there is a window in the overlapping 02 and 03 absorptions from ~ 185 to 210 nm, that is, a region where the total light absorption is in a shallow minimum. This is a region in which the CFCs also absorb light (Fig. 12.14).

The C-Cl bond dissociation energy in CF2C12 is 76 kcal mol-1, whereas that for the strong C-F bond is 110 kcal mol-1. As a result, the weaker C-Cl bond can break at longer wavelengths:

Subsequent reactions of the CF2C1 radical also release the chlorine atoms tied up in this fragment, so that all of the chlorine in the original molecule becomes available for ozone destruction.

Fluorine chemistry in the stratosphere was also considered and it was concluded that ozone depletion by chlorine was > 104 more efficient than that by fluorine (Rowland and Molina, 1975; Stolarksi and Rundel, 1975). Since then, the kinetics of reaction of F atoms with 02 to form the F02 radical and its thermal decomposition have been measured (e.g., see Pagsberg et al., 1987; Lyman and Holland, 1988; Ellerman et al., 1994; and review in DeMore et al., 1997). The equilibrium constant for the F-F02 system

is given by KCI] = 3.2 X 10"25exp(6f00/D (DeMore et al., 1997), so that under stratospheric conditions it lies far to the right, with [F02]/[F] ~104. Potential reactions of F02 that could lead to the destruction of

Wavelength (nm)

FIGURE 12.14 Semilogarithmic plot of the absorption cross sections of the halogenated methanes at 298K: *, CHCl,; ■. CHC1F2; □ , CH2CI2; •. CH2C1F; ▲, CCI4; CCI3F (CFC-11); O, CCI2F2 (CFC-12); CC1F, (adapted from Hubrich and Stuhl, 1980).

Wavelength (nm)

FIGURE 12.14 Semilogarithmic plot of the absorption cross sections of the halogenated methanes at 298K: *, CHCl,; ■. CHC1F2; □ , CH2CI2; •. CH2C1F; ▲, CCI4; CCI3F (CFC-11); O, CCI2F2 (CFC-12); CC1F, (adapted from Hubrich and Stuhl, 1980).

ozone through such cycles as

k%*K < 3.4 X 10"16 cm3 molecule"1 s"1 (DeMore et al., 1997), FO + 03 -» F02 + 02, (23)

< 1 X 10"14 cm3 molecule"1 s"1 (DeMore et al., 1997), have been examined by determining the kinetics of these reactions as well as those of F02 with NO, N02, and the organics CH4 and C2H6 (Sehested et al., 1994;

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