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Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 LECTURE 7 - OZONE LAYER: EXISTENCE AND ANTHROPOGENIC DEPLETION Note: Slide numbers refer to the PowerPoint presentation which accompanies the lecture. Ozone Layer, slide 1 here THE OZONE LAYER The stratosphere has a layer of relatively concentrated ozone at a height of about 16-18 kilometers in the polar regions to about 25 kilometers near the equator. The absorption of ultraviolet radiation by this layer heats the stratosphere, which creates a steep temperature inversion (temperature increases with altitude) between altitudes of 15 and 50 kilometers. This inversion of temperature affects global atmospheric circulation that influences weather and climate. Ozone Layer, slide 2 here Another major effect of the ozone layer is the regulation of ultraviolet light levels at the earth's surface. Ultraviolet light affects all organisms that live on the surface of the land or near the surface of the sea. Marine phytoplankton, marine plants that live near the surface, may be one the most important species affected. Marine phytoplankton provide food for nearly all marine fish, directly or indirectly. Any decrease in the marine phytoplankton will be mirrored by a decrease in marine fish populations. Such a decrease will have severe consequences for marine mammals, for human populations that depend on the sea for food, and for other animals, such as bears, which get part of their food supply from fish returning to freshwater rivers to spawn. Ozone Layer, slide 3 here Ultraviolet light occupies the part of the electromagnetic spectrum from 100 to 400 nanometers (nm). According to the American National Standards Institute (2005), the UV part of the spectrum is divided into UV-A (315 to 400 nm), UV-B (280-315 nm), and UV-C (100-280 nm). Ozone Layer, slide 4 here Different wavelengths are absorbed to different degrees. As the figure shows, the most dangerous type of UV, UV-C, is completely absorbed in the stratosphere. Most, but not all, of the UV-B is absorbed. UV-A is attenuated, but much reaches the ground. Ozone Layer, slide 5 here 1 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 Below 290 nm, absorption of solar radiation is extremely strong. The figure shows the amount of solar energy in watts falling perpendicularly on a surface one square centimeter, and the units are watts per cm2 per nm. The solar energy flux is plotted versus wavelength for four different regions in the atmosphere - the surface, at 20 and 30 kilometers above the surface, and at the “top” of the atmosphere, above 100 kilometers altitude. Below 290 nm in wavelength, the intensity of solar radiation is attenuated by a factor of 108 or more. The blue line, labeled “DNA Action Spectra”, is, “a measure of the relative effectiveness of radiation in generating a certain biological response over a range of wavelengths. This response might be erythema (sunburn), changes in plant growth, or changes in molecular DNA.” (Newman, date unknown) As the figure shows, where DNA damage is likely, the ozone layer in the stratosphere strongly absorbs UV light. The red line in the figure shows the calculated UV spectrum if the ozone levels were reduced 10%, which would lead to an increase of about 22% in DNA damage. Ozone Layer, slide 6 here Increased ultraviolet light levels modify the rate of photosynthesis. Marine phytoplankton are susceptible to such changes. (Goudie, 1990) Crop damage on land is also possible (Stolarski, 1988). Loss of marine phytoplankton could also affect global warming. Phytoplankton absorb carbon dioxide and convert it to oxygen, in the same manner as terrestrial plants. More than half of all the carbon dioxide emitted each year is consumed by plants. Of that amount, more than half is used by phytoplankton. Thus, the phytoplankton are responsible for removal of at least 25% of the carbon dioxide emitted each year. Any reduction in phytoplankton will result in increased carbon dioxide levels in the atmosphere (Vogel, 1993). In 1990 a study of marine phytoplankton near Antarctica by Ray Smith and colleagues showed that phytoplankton losses were about 6-12% in areas exposed to increased UV. Since the ozone hole lasts about three to four months, the annual losses are more like 2-4% (Vogel, 1993). There were substantial differences between species. Diatoms, with siliceous shells that help to protect them, suffered much less than Phaeocystis, a colonial animal growing encased in jellylike material. This can lead to differences in the entire ecosystem supported by the phytoplankton. Ozone Layer, slide 7 here Another potential effect is a change in the emission of dimethylsulfide (DMS). Some phytoplankton manufacture this substance as a kind of natural antifreeze. When they die, the DMS is emitted into the atmosphere. In the atmosphere, DMS acts as a nucleus for cloud condensation. Clouds help to shield UV, so any reduction in DMS emission could lead to fewer clouds and a possible increase in UV (Vogel, 1993). However, this is just speculation. No one knows if the increased UV favors species that make DMS or not. If increased UV favored species that emit DMS, it could act as a negative feedback, increasing cloud cover and decreasing UV. Many such feedbacks exist, but work is needed to gather data. 2 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 Ozone Layer, slide 8 here Foraminiferans, or forams, are single-celled animals who armor their bodies with calcium carbonate shells. In Hawaii, about 25% of the sand on the beach consists of foram shells. Forams sometimes harbor within their bodies photosynthetic organisms capable of providing energy derived from the sun. This is very similar to the process that occurs in corals. Ozone Layer, slide 9 here Since 1987, corals around the world have been ”bleaching.” That is, they suddenly lose their pigment. Now the same phenomenon has been observed in forams. In both cases, the cause is the same. Bleaching is caused by the loss of the photosynthetic organisms. One Florida foraminifera, Amphistegina gibbosa, is now at populations as low as 10% of previous levels (Mirsky, 1994). Remaining animals show developmental defects and malformed shells. High water temperatures have been linked to coral bleaching. Pamela Hollock, of the University of South Florida, investigated whether high-water temperatures were correlated with foram damage. They were not. She found that the damage was strongly correlated with UV light levels at the earth’s surface. She confirmed in the laboratory that UV light levels could cause the kind of bleaching observed in the Florida Keys. The effect may have been triggered by the eruption of Mount Pinatubo, which reduced atmospheric ozone levels and injected dust into the atmosphere. The reduced ozone increased UV levels at the surface, while the dust reduced visible light the organisms could sense. Thinking they were not getting enough light may have made them seek sun, and “fry” themselves (Mirsky, 1994). Ozone Layer, slide 10 here Increased levels of UV-B (UV-B; 280-315 nm) radiation has deleterious effects on living organisms, such as DNA damage (Rozema et al., 1997; Rousseaux, et al., 1999) When exposed to elevated ultraviolet-B radiation, plants display a wide variety of physiological and morphological responses characterized as acclimation and adaptation (Jansen et al., 1998) Ozone Layer, slide 11 here Exposure to different levels of elevated UV-B radiation induced morphological changes and markedly affected the somatic recombination frequency in all the plant lines tested (Ries et al., 2000) To humans, increased ultraviolet levels will translate into increasing rates of skin cancer, including the most deadly form, melanoma. Other effects include an increase in the number of cataracts, and immune deficiencies. Ozone Layer, slide 12 here 3 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 Changes in the ozone layer became of concern in the 1970's. Research on the ozone layer was spurred by perceived anthropogenic threats to the layer. From 1977 to 1984 the ozone levels over Halley Bay, Antarctica decreased by 40% in the region from 12 to 24 kilometers. The problem grew progressively worse into the early 1990's, and the area with the worst ozone depletion is in the Southern Hemisphere. There had been frequent talk of a "hole" in the ozone layer over parts of the Southern Hemisphere, and this hole has grew increasingly large for many years. There have been some signs that a hole may develop over the northern hemisphere. These "holes" are areas of severe ozone depletion, not areas where the ozone layer has disappeared entirely. Over Antarctica the total ozone depletion exceeds 50%. The cause of the ozone hole(s) has been the subject of intense research and debate since the 1970's. It has been suggested that several factors might cause ozone depletion. Some of these are anthropogenic, while others are natural. There is some indication that natural effects are working with anthropogenic effects to significantly lower the ozone levels in the stratosphere. Ozone Layer, slide 13 here The ER-2, a civilian model of the U-2 spy plane, has been used to study the stratosphere. (NOAA/CDML Internet Data). The ER-2 flights are part of the Airborne Arctic Stratospheric Expedition. The ER-2 differs from the U.S. Air Force U-2 in the lack of defensive systems, absence of classified electronics, completely different electrical wiring to support NASA sensors, and a different paint scheme. It also is 30% larger, has 20 ft greater wingspan, and supports a considerably larger payload than the older airframe. To date, NASA U-2 and ER-2 aircraft have flown more than 4,000 data missions and test flights in support of scientific research. (NOAA Climate Monitoring and Diagnostics Laboratory, ER-2, date unknown) Ozone Layer, slide 14 here Other evidence comes from NASA's Upper Atmosphere Research Satellite (UARS). It now appears that significant depletion of ozone could occur in the mid-latitudes and, in a complete surprise to atmospheric scientists, in the tropics. Ozone Layer, slide 15 here In 1992, it was suggested that a hole does might develop over the Arctic, due to findings that suggested conditions favorable for substantial ozone reductions existed over northern Europe, roughly from London to Moscow (Kerr, 1992). At that time, it was suggested ozone losses might reach a depletion of 30-40%. Vogel (1993) stated that harmful ultraviolet striking the northern hemisphere rose 5% in the preceding decade. Over Toronto, Canada, there is evidence of increasing UV-B radiation. It is the most energetic, and most dangerous, form of ultraviolet radiation. Kerr and McElroy (1993) report that ozone levels have decreased 4.1% per year during the winter (December - March) and 1.8% per year during the summer (May to August) between 1989 and 1993. Ninety-three percent of the observed change occurred in the lower 4 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 stratosphere. UV-B measurements correlate with this change. At 300 nm (UV-B), the radiation levels increase 35% per year in the winter and 6.7% in the summer. At 324 nm, the trend was 0.4% per year in the winter and -0.1% per year in the summer. Ozone is strongly implicated because ozone has almost no absorption at 324 nm, whereas it strongly absorbs at 300 nm. Ozone Layer, slides 16-17 here In a 2000 study, it was found that, in one of the Arctic stratosphere's coldest winters on record, measured ozone losses were as high as 55 percent at about 60,000 feet altitude in the ozone layer. The findings may be an indication that future cold winters in the Arctic could prolong the depletion of ozone by manmade chlorine compounds, despite the fact that chlorine is now diminishing in the atmosphere in response to international agreements. (NOAA Climate Monitoring and Diagnostics Laboratory, Large Ozone Depletion Reoccurs this Year in the Arctic, 2000) In answer to the question “Is there depletion of the Arctic ozone layer?” in a 2006 report it was stated: “Yes, significant depletion of the Arctic ozone layer now occurs in some years in the late winter/early spring period (January-April). However, the maximum depletion is less severe than that observed in the Antarctic and is more variable from year to year. A large and recurrent “ozone hole,” as found in the Antarctic stratosphere, does not occur in the Arctic.” (NOAA et al., 2006 Update) Ozone Layer, slide 18 here The results of Kerr and McElroy appear to demonstrate directly that the loss of ozone is responsible for the increase in UV-B radiation (Appenzeller, 1993). Changes during the winter months are large fractional increases in small values. Biologically, they may not be of great significance. Increases during the spring and summer, although of smaller degree, may be more significant. This increase is occurring when many species are reproducing and the biological effects may therefore be larger. The problem will not get better soon. Levels of stratospheric chlorine were about 3.4 ppb in 1990, with a predicted rise to 4.1 ppb by the year 2000, even if all new additions of ozone depleting chemicals to the atmosphere are stopped by the year 2000. Until late 1991 it was thought that further damage to the ozone layer during the 1990's would be comparable to losses experienced during the 1980's. More recent work suggests that the atmosphere's ability to minimize ozone losses is less than was thought. It remains to be seen whether future losses will take the form of scattered spot depletions of ozone within the nonpolar regions, or whether a complete shredding of the stratospheric ozone will occur. In recent years, opinion on the future of the ozone layer has see-sawed back and forth. In a May 2002 article in Nature, Virginia Gewin examined current thinking about ozone and suggested the ozone hole might be repaired by 2040. 5 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 In an excellent commentary in Nature, Dr. Susan Solomon, of the NOAA Aeronomy Laboratory, has made clear that the ozone hole is a problem that is far from cured, and will take considerable time to repair. She warns both scientists and journalists against jumping to conclusions based on yearly data, since so many variable affect the atmosphere. Ozone Layer, slide 19 here The figure shows satellite maps of total ozone over Antarctica on September 24, when the ozone hole is near its annual peak, in 1980, 1981, 2000, 2001, 2002 and 2003. The color scale shows the amount of ozone in Dobson units, indicating the depth of the hole. The images are based on multiple satellite records and analyses. (Solomon, 2004) Ozone Layer, slide 20 here The slide shows the area of the ozone hole in recent years compared to longer term averages, and an image of the current year September 25 ozone hole. Solomon’s warning is well taken. In April and May of 2005, Nature News articles reported that the Arctic ozone hole suffered the biggest losses ever reported, and that the protective ozone layer over northern and central Europe was thinner in 2005 than it had been since measurements began 50 years previously. Scientists were arguing whether global warming was to blame. The argument is that increasing levels of greenhouse gases have a cooling effect on the stratosphere, because heat is locked near the surface. As we shall see, it takes very cold temperatures in the stratosphere to produce an ozone hole. Once again, Susan Solomon urged caution because the temperature drop during the 20042005 winter was far too large to be explained by greenhouse cooling. She added that natural fluctuations in meteorological conditions probably far outweigh greenhouse effects. (Schiermeier, 2005a and 2005b) Current theories on ozone depletion will be examined. We will first examine anthropogenic causes, then look at natural contributions and possible synergistic effects of the natural and anthropogenic causes that together might pose a greater danger than either alone. ANTHROPOGENIC OZONE DEPLETION Ozone Layer, slide 21 here Serious changes in the ozone levels in the stratosphere began to be investigated in the 1970's. At first, it was thought that fleets of supersonic transport planes (SST's) would introduce water vapor and nitrous oxides into the stratosphere, degrading the atmosphere. Fleets of these aircraft have not materialized. Ozone Layer, slide 22 here In 1974, another threat was recognized. This threat was the use of chlorofluoro-carbons (CFC's) (Molina and Rowland, 1974). As previously discussed, from 1977 to 1984 the ozone levels over Halley Bay, Antarctica decreased by 40% in the region from 12 to 24 kilometers. Ozone Layer, slide 23 here 6 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 Molina and Rowland, together with Paul Creuzen, shared the 1995 Nobel Prize for Chemistry for the timely work and warnings about the ozone hole problem. Ozone Layer, slide 24 here In the Antarctic the ozone hole typically begins developing in August, reaches a maximum in early October, and disappears by early December. The cause of this ozone depletion has been attributed to several anthropogenic causes. These include: Ozone Layer, slide 25 here 1. Combustion products from high-flying military and civilian aircraft, particularly supersonic aircraft 2. Nitrous oxides released from nitrogenous fertilizers 3. Chlorofluorocarbons (CFC's), first introduced in the late 1920's, are used as refrigerants, in the manufacture of foam fast-food containers, as cleansers for electronic parts, and as propellants in aerosol cans 4. Other compounds, such as Halon and methyl bromide, which contain bromide, and which are capable or releasing substances capable of destroying ozone. Of the four, the CFC's might be the most serious problem, and have received the most attention from scientific researchers and from the press. However, as the CFC problem has been addressed, the bromide compounds assume larger relative importance. Therefore, we shall also examine these compounds and their effect on the ozone layer. Ozone Layer, slide 26 here There are many common CFC's. Some are called Freon. Freon is not a single compound. There are various Freons, usually designated by number (see Table 7-1). Freons are often used as refrigerant gases. Today, HFC or HCFC compounds are being substituted for Freon. Halon gases are used in fire extinguishers. This use has now been largely banned. All CFC compounds have chlorine and/or fluorine atoms substituted for hydrogen. The Halon compounds also have bromine, which causes its own problems 7 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 Table 7-1 Common Chlorofluorocarbons Normal Designation, Other name Formula Chemical name CFC-11, Freon 11 CFCl3 Trichlorofluoromethane CFC-12, Freon 12 CF2Cl2 Dichloro-difluoromethane CFC-13, Freon 13 CF3Cl Chloro-trifluoromethane CFC-113 C2F3Cl3 Trichloro-trifluoroethane CFC-114 C2F4Cl2 Dichloro-tetrafluoroethane CFC-115 C2F5Cl Chloro-pentafluoroethane HCFC-22 CHF2Cl Chloro-difluoromethane HCFC-123 CHCl2CF3 Dichloro-trifluoroethane HCFC-124 CHFClCF3 Chloro-trifluoroethane HFC-125 CHF2CF3 Pentafluoroethane HCFC-132b C2H2F2Cl2 Dichloro-difluoroethane HFC-134a CH2FCF3 Tetrafluoroethane HCFC-141b CH3CFCl2 Dichlorofluoroethane HCFC-142b CH3CF2Cl Chloro-difluoroethane HFC-143a CH3CF3 Trifluoroethane HFC-152a CH3CHF2 Difluoroethane HALON 1211 CF2BrCl Bromochloro-difluoromethane HALON 1301 CF3Br Bromo-trifluoroethane HALON 2402 C2F4Br2 Dibromo-tetrafluoroethane After Houghton et al., 1990, Appendix 8 8 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 Ozone Layer, slide 27 here The bonds between carbon and chlorine, carbon and fluorine, or carbon and bromine, are stronger than the bonds between carbon and hydrogen (Table 7-2). Table 7-2 Relative bond strengths Bond Bond Strength, kcal/mole C-H 80.9 C-Br 95.6 C-F 107 Data from Weast, 1966 Thus, CFC compounds are very stable. They do not break up in the troposphere, of the earth's atmosphere. They circulate without change in the troposphere and some eventually rise into the stratosphere. Sherwood Rowland recounted his work with then graduate student Mario Molina in elucidating the removal mechanisms of ozone in the atmosphere. He said in his Nobel Prize lecture, (Rowland, 2007) Ozone Layer, slide 28 here “When Mario Molina joined my research group as a postdoctoral research associate later in 1973, he elected the chlorofluorocarbon problem among several offered to him, and we began the scientific search for the ultimate fate of such molecules. At the time, neither of us had any significant experience in treating chemical problems of the atmosphere, and each of us was now operating well away from our previous areas of expertise. Ozone Layer, slide 29 here 9 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 The search for any removal process which might affect CCl3F began with the reactions which normally affect molecules released to the atmosphere at the surface of the Earth. Several well-established tropospheric sinks - chemical or physical removal processes in the lower atmosphere - exist for most molecules released at ground level: Ozone Layer, slide 30 here 1. Colored species such as the green molecular chlorine, Cl2, absorb visible solar radiation, and break apart, or photodissociate, into individual atoms as the consequence; 2. Highly polar molecules, such as hydrogen chloride, HCl, dissolve in raindrops to form hydrochloric acid, and are removed when the drops actually fall; and Ozone Layer, slide 31 here 3. Almost all compounds containing carbon-hydrogen bonds, for example CH3Cl, are oxidized in our oxygen-rich atmosphere, often by hydroxyl radical as in reaction (1). CH3Cl + HO H2O + CH2Cl (1) Ozone Layer, slide 32 here However, CCl3F and the other chlorofluorocarbons such as CCl2F2 and CCl2FCClF2* are transparent to visible solar radiation and those wavelengths or ultraviolet (UV) radiation which penetrate to the lower atmosphere, are basically insoluble in water, and do not react with HO, O2, O3, or other oxidizing agents in the lower atmosphere. When all of these usual decomposition routes are closed, what happens to such survivor molecules?” Ozone Layer, slide 33 here 10 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 Their stability was at first thought to make them an ideal industrial compounds because they are unreactive and therefore nontoxic. Carbonchlorine bonds can be broken by ultraviolet radiation, releasing chlorine free radicals (equation 7-1). Ozone Layer, slide 34 here 7-1 The chlorine free radical then reacts with ozone: 7-2 Chlorine monoxide can further react, 7-3 or, 7-4 Chlorine atoms thus released by equations 7-3 or 7-4 can again react according to equation 7-2. The net result is a chain reaction that destroys thousands of ozone atoms for each chlorine atom initially released. Chlorine is acting as a catalyst, since it is not consumed in the total reactions. It is the magnification of the chain reaction process that makes the CFC's so dangerous. Eventually the chlorine atom may migrate back to the troposphere. Here the chlorine will react to form hydrochloric acid, which will be removed by rain (and thus become a minor contributor to acid rain). When the CFC compounds reach 25 kilometers or more altitude, the ultraviolet light is strong enough to initiate bond breakage. The ozone levels above 25 kilometers are small and thus the ultraviolet levels are higher. The release of chlorine triggers the rapid depletion of ozone because of the chain reaction. 11 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 Even if all CFC emission were to stop immediately, the problem would not disappear. Freon 11 persists for about 75 years and Freon 12 for about 100 years in the atmosphere (Stolarski, 1988). Since most of the CFC's are in the troposphere, and are only slowly migrating into the stratosphere, the stratospheric CFC levels will increase for years. Ozone Layer, slide 35 here Freon 11 (Elkins 2016) Ozone Layer, slide 36 here Freon 12 (Elkins 2016) Monthly means of the dry mixing ratios in parts per trillion of Freon 11 and Freon 12 are shown. Ozone Layer, slide 37 here HCFC Plots of HCFC data for three compounds. Note the scales. HCFC-22 is about 10 times as much as HCFC-141b or HCFC-142b. Ozone Layer, slide 38 here Halons and methyl bromide Plots for two halon compounds are shown. Halon compounds contain bromine, and are often used as fire suppressants because the vapor is heavy. Methyl bromide was heavily used as a soil fumigant for killing and controlling a variety of pests. Because the vapor is heavy, it has low volatility and is retained by soils for long periods. Ozone Layer, slide 39 here There have been attempts to control CFC use. In 1978, the United States banned the use of CFC's in aerosol products such as hair sprays and deodorants. Britain, in 1984, announced their data showing 40% ozone depletion. In September, 1987 23 nations, including the United States, signed an agreement, known as the Montreal Protocol, to reduce consumption of chlorofluoro-carbons (United Nations Environment Programme, 2000). It required the signatory nations to reduce consumption of CFC's to 1986 levels 12 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 by the mid-1990's and to halve the usage by 1999. Methyl bromide produced and imported in the U.S. was reduced incrementally until it was phased out in January 1, 2005, pursuant to our obligations under the Montreal Protocol on Substances that Deplete the Ozone Layer (Protocol) and the Clean Air Act (CAA). (EPA Ozone Layer Protection, 2014) Ozone Layer, slide 40 here Chlorine loading is a measure of the real contribution of a substance to ozone depletion. A projection of future halocarbon contributions to chlorine loading is shown in the figure. All ozone depleting substances are shown on a common scale that takes account of differences in potency. Bromine is expressed as its equivalent in chlorine. (European Fluorocarbons Technical Committee, 2006). Fortunately there are reactions that destroy chlorine in the stratosphere. An important degradation reaction involves either chlorine atoms or chlorine monoxide molecules. These react with another molecule to form a "stable" product that temporarily removes the chlorine into a chlorine reservoir. Chlorine trapped in the reservoir cannot attack ozone. Ozone Layer, slide 41 here Two important reactions are the reaction of chlorine monoxide and nitrogen dioxide to form chlorine nitrate, 7-5 and the reaction of chlorine and methane to form hydrochloric acid: 7-6 The CH3 radical is reactive and will undergo further reactions. Unfortunately the molecules in the reservoir will absorb a photon and break apart, releasing chlorine. Thus these reactions slow the destruction of ozone, but probably prolong the destruction over many years. It is ironic that the various nitrogen 13 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 oxides are themselves thought to destroy ozone, but here serve to break the chlorine chain reaction and thus conserve ozone. Ozone Layer, slide 42 here Nitrogen gases attack ozone approximately as follows. Nitrous oxide (N2O) is produced as a by-product of nitrification and denitrification in soils and oceans. In the troposphere, there are no known degradation reactions for nitrous oxide, so it migrates upward into the stratosphere. Most of the nitrous oxide will be destroyed by photodissociation into oxygen atoms and N2. Approximately 10% escapes this fate and is destroyed by reaction with activated atomic oxygen. This reaction produces either nitric oxide (NO) or N2 and O2. It is the nitric oxide that is important because of the ability of NO to catalytically destroy ozone (Kinzig and Socolow, 1994). Nitric oxide can also be directly injected into the stratosphere by high-flying supersonic planes. This was originally thought to be the biggest danger to the ozone layer. Ozone Layer, slide 43 here 7-7 7-8 This yields a net reaction of: 7-9 Nitrous oxide additions to the atmosphere can be divided into natural and anthropogenic components. Natural additions, from temperature and tropical soils and the world’s oceans are estimated at 10 million tons per year. Anthropogenic emissions are thought to be of similar magnitude. The most important source is agricultural use of nitrogen fertilizers. Other sources include industry (production of fertilizer and nylon), fossil fuel use in power production, biomass burning, cattle and cattle feed production. (Reay, 2007) 14 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 POLAR STRATOSPHERIC CLOUDS AND HETEROGENEOUS CHEMICAL REACTIONS Ozone Layer, slide 44 here The presence of the inhibitor reactions led the computer models to predict that CFC's should have had little effect on the ozone layer by the late 1980's. Yet the ozone hole over Antarctica was unmistakably present, so something was wrong. Ozone Layer, slide 45 here Antarctica is different from the temperate parts of the globe in many ways. One important way is the presence of polar stratospheric clouds, or PSC's. These clouds form during the Antarctica winter, when the absence of the sun for long periods lets the stratospheric temperatures dip below -78C (<195 Kelvins). Ozone Layer, slide 46 here The figure in the PowerPoint presentation shows the temperatures over the Neumayer station, at 70 south latitude, during 1997. It has been known for most of the last century that stratospheric clouds occasionally formed over both the north and south pole. Ozone Layer, slide 47 here - Polar Stratospheric Cloud Photos They often glow with seashell-like iridescence. This has led to the name nacreous, or mother-of-pearl, clouds. Besides the nacreous clouds (called Type II), composed of water ice, two other types exist. Ozone Layer, slide 48 here The video shows nacreous clouds. Description by Laura Candler, the photographer: “On January 11, 2010, a beautiful group of nacreous clouds 15 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 appeared over Kiruna, Sweden. Also known as mother-of-pearl clouds, nacreous clouds sometimes form in the polar stratosphere in wintertime and glow with a silky iridescence as they undulate across the sky. I created this time lapse using images shot at 10 second intervals over the course of about 2 hours, from 10:25 a.m. until 12:14 p.m.. The images have not been manipulated in any way to enhance color, exposure, “ Ozone Layer, slide 49 here One consists of nitric acid ice (Type I). The third type is similar in chemistry to the nacreous clouds, but the cloud is larger and shows no iridescence (Type III). In 1986, Susan Solomon and colleagues and Michael McElroy and coworkers provided the first hint as to the connection between PSC's and ozone depletion. They suggested that chemical reactions on the surface of ice crystals in the clouds release chlorine. This chlorine in turn reacts with thousands of ozone molecules before again being captured into a reservoir. Reactions on the surface of particles are called heterogeneous chemical reactions. This suggestion spurred research on the PSC's. Ozone Layer, slide 50 here The Type I cloud, containing nitric acid, is usually hard to observe from the ground. The Stratospheric Aerosol Measurement (SAM) II, launched on the NIMBUS satellite in 1978, detected particles in the air by examining sunlight grazing the earth and passing through the stratosphere. SAM II indicated the presence of clouds even when the temperature was 195 Kelvins, too warm for water ice to form. It was proposed in 1986 that these clouds consist of nitric acid trihydrate (NAT =HNO3.3 H2O) (Toon et al., 1986). At these temperatures, which are above the freezing point of water ice under ambient conditions, the nitrogen compounds may freeze onto aerosol particles (probably sulfuric acid aerosol), becoming part of the ice in the clouds. This would remove the inhibitor molecules, allowing more chlorine atoms to form. Such clouds are produced by slow cooling, due to radiation cooling in the long polar nights. Consequently, they are more common in the Antarctic than the Arctic, which experiences more mixing with the warmer 16 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 mid-latitude air. The nitric acid clouds are more tenuous than nacreous clouds, but are often much larger, reaching sizes up to several thousand kilometers. Ozone Layer, slide 51 here The ice particles also serve as catalysts for the destruction of chlorine reservoir molecules. Laboratory studies showed that a reaction between hydrochloric acid and chlorine nitrate (HCl and ClONO2) produces photolytically unstable species such as Cl2, ClNO2, or HOCl on the surface of ice crystals (Toon and Turco, 1991). Sunlight transforms these species into the ozone-reactive species Cl or ClO (Cicerone et al. 1991). Reactions often take place much faster on a solid surface than in a purely gaseous phase. Without ice, the reactions are negligibly slow. Overall, the lack of sunlight would be expected to slow most chemical reactions to near zero levels. However, these processes will release chlorine or chlorine monoxide quickly when the sun returns. This would lead to a rapid ozone depletion in the stratosphere above Antarctica at the start of the southern spring, in agreement with what has been observed. Ozone Layer, slide 52 here The nacreous clouds look like lenticular clouds which form in mountainous regions. However, lenticular clouds form in the troposphere. In the stratosphere, the relative humidity is about 1%, and the average water vapor content is a few parts per million. It is difficult to imagine ice clouds forming under these conditions. The lenticular clouds are known to form as air rushes up a mountain. Such an air mass undergoes rapid cooling, increasing the relative humidity. This can create a series of clouds in a standing wave pattern downwind from the mountain. Under certain conditions, the standing wave pattern can propagate upward into the stratosphere. Nacreous clouds are thought to form by condensation of ice on any aerosols present on the crests of the standing waves. Air is constantly passing through the nacreous clouds, but the clouds are stationary. At the leading edge, ice begins to nucleate to form a cloud. The ice particles grow 17 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 to about 2 microns in size before exhausting their supply of water. As the crystals move with the air they eventually reach the descending portion of the airmass, on the lee side. As they descend, they sublimate. The cloud disappears. Thus the clouds appear stationary, although the air and ice crystals are constantly moving. Nacreous clouds are important because they tell us that temperatures are below 190 Kelvin, the temperature at which ice crystallizes under conditions in the stratosphere. (Toon and Turco, 1991) The iridescence observed in nacreous clouds is due to their size and the density of the ice particles in the clouds. As light passes through the cloud it passes through many crystals, and is refracted in each. Since different wavelengths of light are refracted at different angles and absorbed differently, iridescence is produced. The effect is similar to iridescence in minerals. Ozone Layer, slide 53 here Work by Webster et al. (1993) has shown that the reaction, 7-10 occurs in Type I or II PSC’s at temperatures below 196 ± 4K on particle surfaces in the clouds. This reaction begins with an excess of ClONO2, not HCl, as was earlier believed. It is this reaction that appears responsible for partitioning of chlorine between various phases. This work was based on actual in situ measurements of species in the atmosphere. Toohey et al. (1993) have shown increased ClO levels in the Arctic polar vortex during the 1991-92 winter when temperatures descended below 195 K. They also observed rapid decrease in ClO concentrations in February 1992 as temperatures warmed and ClONO2 formed. Profitt et al. (1993) showed that there was an altitude dependent change in ozone concentration in the Arctic polar vortex. There was a substantial decrease toward the bottom of the vortex, which they attributed to ozone-rich air entering the top of the vortex. The ozone-rich air reacts with high concentrations of reactive chlorine within the cloud. The ozone-depleted air released from the bottom of the vortex is of sufficient quantity to significantly lower the mid-latitude ozone levels. 18 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 Ozone Layer, slide 54 here Type III polar stratospheric clouds are similar to the nacreous clouds, except that the particle size is much larger, about ten microns. Temperature must drop slowly below 190 Kelvin. Water ice begins to form on seed particles, either nitric acid particles or any remaining sulfuric acid aerosol. Cooling is slow and the particles can grow to large size. The particle density is much less than in a nacreous cloud. Since the particle size and density are different from that of nacreous clouds, these clouds are not iridescent. Slowly-cooled water ice clouds are important because they remove nitrogen from the atmosphere. They form on nitric acid substrate and remove nitric acid vapors from the air. This decreases the nitrogen available for complexing the chlorine into a reservoir. There is one additional problem that must be overcome during the southern spring. The sun is low on the horizon, which reduces the radiation driven breakdown of ozone. This, in turn, reduces the number of oxygen atoms available for the chlorine catalytic cycle of ozone destruction (equation 7-4). It has been suggested (Stolarski, 1988) that the presence of bromine in the stratosphere could help overcome the lack of oxygen atoms. Bromine comes from several sources. One is naturally occurring methyl bromide. Anthropogenic sources include fumigants and certain fire extinguishers. Ozone Layer, slide 55 here Bromine can interact with ozone as follows: 7-11 The bromine monoxide can interact with chlorine monoxide to form a bromine atom plus a chlorine atom. 7-12 The net result is the conversion of ozone to oxygen. Thus, a brominechlorine cycle could operate even lacking free oxygen atoms. Br can also 19 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 destroy ozone via chain reactions similar to the chlorine catalytic reaction, but not requiring free oxygen atoms. Methyl bromide is now regulated, and has been eliminated by law (see Appendix A). Ozone Layer, slide 56 here There is some data to support the ideas outlined above. Data from a major study in 1986 at McMurdo Sound, Antarctica, and other work indicate that the springtime levels of chlorine monoxide are elevated over Antarctica relative to mid-latitude stratospheric sites, and that nitrous oxide levels are severely depressed. Also, both chlorine nitrate and hydrochloric acid levels are low simultaneously. Work in the Arctic during the winter of 1991-92 showed that the levels of ClO increase in December, when the stratosphere first gets cold enough for PSC’s to form. ClO concentration reaches a maximum in January, when temperatures are cold enough that most of the air inside the polar vortex has been processed inside PSC’s. In February, ClO concentration begins to decrease because air temperatures are warming. Without the huge polar icecap, the Arctic temperatures warm much earlier than the Antarctica temperatures. Thus most of the ClO is gone before sufficient sunlight is available for catalytic ozone removal to occur. In Antarctica, the ozone removal reaches a maximum in early to mid-October, which is equivalent to April in the northern hemisphere (Rodriguez, 1993). HETEROGENEOUS REACTIONS OUTSIDE POLAR VORTICES Ozone Layer, slide 57 here Heterogeneous reactions have also been observed outside polar stratospheric clouds. Hydrochloric acid molecules have been expected to be the major component of inorganic chlorine in the atmosphere. Normally, hydrochloric acid is expected to be very stable and inactive. Reactions between HCl and ClNO3 have been shown to occur rapidly on PSC -like laboratory surfaces. HCl values do indeed decrease in the PSC’s. However, 20 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 measurements of HCl outside polar vertices also showed that HCl was sometimes lower than predicted. This means that HCl may not be the dominant species, and raises the question of what unexpected reactions are occurring. The results of the second Airborne Arctic Stratospheric Expedition (AASE II) have provided some possible answers. Ozone Layer, slide 58 here After the eruption of Mount Pinatubo in 1991, there was a substantial increase in global aerosol loading. The spreading volcanic cloud reached a maximum enhancement in particle surface area of thirty times that observed during AASE I in 1989. Airborne instruments can sample and analyze changes in aerosol size and surface area along with monitoring changes in chemistry (Rodriguez, 1993). It was found that a “footprint” of heterogeneous reactions on ClO existed despite the absence of any polar stratospheric clouds in some areas (Wilson et al., 1993). Similar reductions were found in nitrogen oxides. If temperatures are cold enough, the reactions on sulfate aerosol can be as efficient as those within PSC’s. This will be discussed more fully in lecture 8. Ozone Layer, slide 59 here A Dobson unit (DU) is the standard method of measuring ozone levels. The Dobson unit is 0.01mm, or 10 microns. For ozone, the Dobson units represent the thickness of the layer of ozone above a given point if all of the ozone could be brought together at a standard pressure and temperature. ATMOSPHERIC DYNAMICS AND OZONE Ozone Layer, slide 60 here Geochemical degradation may not be the only cause of the ozone hole. The earth's atmosphere is far from static. There is circulation at all levels of the atmosphere. This circulation helps to keep the temperature of the earth more uniform than solar heating levels would otherwise suggest. In the 21 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 stratosphere, ozone is produced most strongly at high altitudes and low latitudes. Thus, we might expect that the greatest concentrations of ozone would be found above the equator at the top of the stratosphere. Actually, the ozone levels peak in the mid-stratosphere and near the poles. This distribution is caused by stratospheric circulation from high altitudes above the equator to lower altitudes near the poles. The ozone is carried along with the air mass. In the northern hemisphere peak ozone levels reach 450 Dobson units in the late winter or early spring, over the North Pole. In the southern hemisphere circulation is limited to about 60 south latitude. There is little exchange between the polar vortex and mid-latitude regions. Peak ozone levels reach 380 Dobson units at 60 south latitude. At the equator, ozone levels are typically about 260 Dobson units (Stolarski, 1988, p.35) Ozone Layer, slide 61 here In the polar vortex above Antarctica the ozone levels once were about 300 Dobson units during the winter and early spring. The amount would increase to about 400 Dobson units in the late spring when the polar vortex broke down, allowing exchange with mid-latitude air masses. Now the ozone levels are steady throughout the winter but fall rapidly in the spring to less than 200 Dobson units. The early winter ozone minimum is a natural phenomenon. Ozone Layer, slide 62 here It is the depletion below 300 Dobson units that represents the ozone hole. In 1991 the record low value of ozone minima first fell below 100 DU over Antarctica, and that has had in ten additional years since then. The record low was September 30, 1994, with the lowest value ever recorded at 73 DU. It was anticipated that the 1994 values would be much higher than the 1993 values, since most of the sulfur gases from the 1991 Mount Pinatubo eruption had washed out of the atmosphere. This was proven not to be the case (Kerr, 1994, and NASA Goddard Space Flight Center, 20130. Ozone Layer, slide 63 here 22 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 Thus atmospheric circulation may contribute to changing ozone levels, but do not appear able to explain the ozone hole observed over Antarctica. CFC's remain the most likely culprit. A plot of the ozone depletion over the South Pole station at the maximum ozone hole formation in 1994 and 2015 shows little difference. Appendix A The following is taken from the Internet (at http://aceis.agr.ca/policy/env/methbrom.html), and was written by Linda Dunn, of the Environment Bureau, Agriculture & Agri-Food Canada. OPTIONS FOR METHYL BROMIDE When Methyl Bromide, used worldwide as a fumigant for pest control in soils, structures, spaces and commodities, was identified as an ozone-depleting substance, it came under the control of the Montreal Protocol. Parties agreed in November 1992 to freeze production and consumption of Methyl Bromide at 1991 levels by 1995, except for quarantine and pre-shipment applications. Parties are required to meet this commitment, however they can also implement more stringent domestic actions. Canada committed to a further 25% reduction by 1998, an action recently matched by the European Union. The freeze and reduction in Canada have been incorporated into changes to the Canadian Environmental Protection Act (CEPA) which came into effect January 1, 1995. The United States will eliminate production and consumption by the year 2001, under their Clean Air Act. As agriculture is the primary user of Methyl Bromide in Canada, it falls on our sector to find ways to reduce use. The Environment Bureau of Agriculture & Agri-Food Canada (AAFC) has been co-chairing the process with Environment Canada. A workshop was held in November 1993 with interested stakeholders. As a follow-up, an informal industry / government working group was established to look at research priorities and possibilities for joint projects, effective ways to implement alternatives and regulations, ways the industry can reduce emissions, and assisting in the development of Canadian positions 23 Environmental Geochemistry, GLY 4241/5243, © David Warburton, 2016 on future controls. The industry has responded to this challenge by designing its own program of tradable permits. Through the creation of a market for these permits, the industry has taken ownership of this issue and the approach to addressing it. Last year's meetings to the Montreal Protocol focused on clearly defining quarantine and pre-shipment exemptions. Canada's goal, which was attained at the 1994 meeting of the Parties, was to push for definitions which closely matched those under CEPA. When Parties to the Protocol placed Methyl Bromide under interim controls in 1992, they requested a number of studies to be undertaken, including one on alternatives. The Plant Industry Directorate of AAFC was Canada's representative on this committee. All of the studies are now completed, and in time for the November 1995 meeting of the Parties. Two working group sessions will be held prior to this meeting, and will make recommendations on further controls (including phase-out dates). A national consultation was held in March, 1995, to discuss issues related to further controls and advise on a Canadian position leading up to the 1995 meetings. 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