ENVR 30, P. Chau Chapter 10

Transcription

1 ENVR 30, P. Chau 2007 Chapter 10 Aiir polllluttiion regiionall and glloball probllems Acid precipitation and the depletion of the stratospheric ozone layer are two air pollution problems that draw regional and global concerns. Problems of acid rain will linger as long as we burn fossil fuels, especially coal, without adequate pollution control. The international community has banned the use of ozone layer depleting chemicals, but it will still be decades before the ozone layer can recover. A brief history of acid rain. An English chemist, Robert Smith, appears to be the first one who observed the phenomenon of acid rain around the city of Manchester back in Another Englishman, Eville Gorham, showed as early as 1955 that much of the acidic precipitation was attributed to combustion emissions near industrial regions. In 1961, Svane Odin, a Swedish soil scientist, noted that acid precipitation was a regional problem and correctly identified that it was due to the transport of sulfur- and nitrogencontaining chemicals over long distances, as much as some 2000 km. Odin also predicted that acid rain would lead to the decline of fish populations, leaching of toxic metals from soil, and decreased forest growth. Concern about acid precipitation in North America developed first in Canada in the 1970s because lakes in southeast Canada are less buffered and thus more susceptible to acid precipitation. Little was done in the U.S. until the 1990 Clean Air Act which aimed to reduce sulfur dioxide emissions in the Midwest power plants by half. What is acid precipitation? Acid precipitation is caused mainly by SO 2 and NOx emissions from the burning of fossil fuels. Atmospheric SO 2 and NOx are converted to sulfuric acid (or sulfates) and nitric acid (or nitrates). These secondary pollutants are carried in tiny liquid or solid particulates. These acidic particulates return to earth in the form of rain, fog, mist, and snow (wet deposition or precipitation) or just dry acid particles (dry deposition). The whole collective process may be considered as acid deposition. On the whole, we have better programs to monitor the ph and chemical composition of rainfall than the other forms of precipitation. Thus, we commonly address the entire phenomenon as acid rain. We also know that natural emissions of SO 2 and NOx can be substantial (see Chapter 9). The atmosphere also has CO 2 which forms a weak acidic solution in rain droplets. So one may question what should be considered as really acid rain. If natural rain droplets are in equilibrium with the global atmospheric CO 2 concentration, this clean raindrop will be slightly acidic with a ph value of 5.6 (see the Introduction chapter for the definition of ph). With probable background of natural sulfur compounds, the ph of the rainfall may drop to about 5.0. Thus raindrops with ph between 5 and 5.6 may have been influenced by humans but not to an extent that exceeds natural background levels. Hence, we usually consider acid rain as precipitation with ph less than 5.0. Indeed, many regions that suffer from acid precipitation (northeastern U.S., southeastern Canada, Mississippi basin) receive rainfall with ph less than 4.5. In California, precipitation with ph between 4 and 5 was once common before the enforcement of air pollution laws.

2 10 2 The source of acid precipitation and controversy. We know that the cause of anthropogenic emission of SO 2 and NOx is predominantly combustion of fossil fuels. Since about two-thirds of the acidity of the precipitate is due to sulfuric acid and about three-fourths of the SO 2 comes from coal-fired power plants, a significant amount of acid precipitation would be attributed to these utility plants. In particular, coal-burning power plants that have no emission control except for a huge smoke stack (500 to 1,000 feet tall) would be a prime culprit for the amount of SO 2 released. Similar situations also exist in many eastern and western European countries. However, the location of the power plants can be hundreds of miles away from the impacted areas. Unfortunately, regulatory agencies have the burden of proof, and a key issue in acid rain is the exact source of emissions that ultimately impacts a given area as acid deposition. The determination of the cause-effect (or source-receptor) relationship involves two important elements: (1) identification of the trajectories of an air current that arrives at the deposition location, and (2) prediction of the chemical composition of the pollutant based on the suspected source. To establish a causal link when the impact area is hundreds of miles away is a formidable task. It took years of measurements, experimentation, and computer modeling before scientists finally established a strong legal case to link SO 2 and NOx emissions to their environmental impacts. Not all lakes are born equal! Presence of limestone provides buffering capacity. One early controversy was that many areas were not affected equally. For example, lakes in southeast Canada are more susceptible to acid rain than lakes in northeast America. The reason lies in the buffering capacity offered by limestone. Alkaline soils that are rich in limestone can neutralize acid. When limestone (CaCO 3 ) is present in the soil or bedrock, it reacts with acid (hydrogen ion H + ) to form bicarbonate ion: H + + CO 2 3 > HCO 3 The bicarbonate ion can further react with acids to form the weak bicarbonate acid (H 2 CO 3 ): H + + HCO 3 > H 2 CO 3 Moreover, bicarbonate ions (HCO 3 ) may also be present in natural waters to help neutralize acids. As long as there is excess limestone, the ph in the aquatic ecosystem is fairly stable, despite acid rain. Hence, the acid-neutralizing capacity of the lake buffers it against large changes in ph. Of course, the size, depth, and the drainage of a lake (geology, vegetation, water flow) also influence the ultimate response of the lake to acid rain. Once the limestone has reacted with acid, it is removed from the soil. As soon as the limestone or the source of bicarbonate ions is depleted, the buffering capacity is lost and the ph now drops precipitously with the addition of acid precipitation. (This is very much like doing titration in your chemistry lab and overshooting the end-point.) Indeed, sudden or rapid collapse of lakes and forests has been observed. Now back to the Northeast region of America, the lakes in Canada are lined with thin soil and glacial granitic bedrock and their buffering capacity can be exhausted quickly. In contrast, the lakes in America tend to have a higher content of limestone in the soil of the watershed

3 10 3 Acid rain can leach metal ions and harm forest growth. The buffering can also be provided by cation exchange, which can be illustrated with the chemical reaction: 2H + + M 2+ -Soil > M H + -Soil Here, M 2+ represents some alkaline metal cation (positive ions such as calcium Ca 2+ and magnesium Mg 2+ ) that is bound to soil particles. The hydrogen ion H + displaces these metal ions from soil minerals. The metal ions are released while hydrogen ions are retained in the soil (a process also called mobilization, or in laymen terms, leaching). When calcium and magnesium, essential for plant growth, are first leached from soil particles, plants can absorb some of them, but rainfall and streams carry the rest away. Soon, the soil is depleted of the nutrients, retarding forest growth and recovery. With the same displacement process, the acid also helps to dissolve aluminum and other heavy toxic metals from rock and soil minerals. It has been observed that as the ph of a lake decreases, concentrations of metals such as aluminum, iron, copper and zinc increase. Presence of toxic metals is a contributing factor in the decline of forest growth. In reality, multiple factors are responsible for dying and damaged forests. Various pollutants work together synergistically to amplify the damage. Environmental impact of acid rain. The immediate effect that one observed as early as the 1950s was the decline in fish populations in lakes and the death of trees in Europe. Here is a summary of the adverse effects of acid rain: Effects on aquatic ecosystems Natural freshwater has a ph around 6-7 and aquatic organisms have adapted to such fairly neutral waters. As the ph drops below their optimal range, they are stressed (yes, back to ecosystems in Chapter 2). The eggs of amphibians and trout are especially sensitive to ph changes. Eggs, sperm, and the young are more severely affected, lowering recruitment. Of course they die if the ph drops below their limit of tolerance, which for most fish is between 4 and 5. An acidified lake, with its acid-neutralizing capacity exhausted, usually has a ph well below 6 and its waters are high in sulfate and metal ions such as aluminum. Aluminum is harmful to plants and aquatic life. Finally, heavy metal leached from soil can cause abnormal development and death of fish embryos. Effects on forests Acid rain can lead to reduced growth and dieback of forests by several means. The mobilization of aluminum and other toxic heavy metals may adversely affect tree growth, as we have explained in the cation exchange section. At the same time, acid rain displaces useful inorganic ions such as calcium and magnesium from soil particles. Plant growth will slow down with the depletion of nutrients. High levels of nitrate can form nitric acid and kill mycorrhizae, a symbiotic fungus that lives on the roots of conifers and help trees to extract water and nutrients. Low ph also retards the activity of the detritus community and the rate of nutrient cycling.

4 10 4 Acid deposition can also damage the protective layer of leaves or needles of conifers. Meanwhile, ozone may attack the trees. As the trees are weakened under environmental stress, they are more susceptible to infection and disease. Hence many factors work simultaneously or synergistically and it is very difficult to isolate one single cause for the demise of forests. Effect on humans Many historic statues and monuments were built of limestone and marble. Acid rain or fog has caused alarming weathering and erosion of these invaluable artifacts. A brief history of ozone layer depletion. Ozone (O 3 ) is a very important trace compound in the stratosphere. Although it is present only in trace amount, ozone is responsible for shielding the earth from ultraviolet radiation that is harmful to life. The ozone layer absorbs much of the sun's ultraviolet radiation, particularly the so-called ultraviolet-b radiation that is especially dangerous to plants and animals. Chlorofluorocarbons (CFCs, was once referred to as freons, their trade name) are a family of chemicals that is very volatile but nontoxic and chemically stable. The key species are in the subclass chlorofluoromethanes: CFCl 3 (CFC-11) which was widely used as a propellant in aerosol spray cans and in making plastic foams, and CF 2 Cl 2 (CFC-12) which was used in air conditioners and refrigerators. These chemicals were first introduced as aerosol propellants during the Second World War. Until the mid-1990s, CFCs were still used for cleaning in the electronic and aerospace industries. Within a few years of their release, the CFCs were distributed quite uniformly throughout the troposphere. Because they are chemically inert in the troposphere, their atmospheric lifetimes are extremely long. For CFC-11 and CFC-12, they are estimated as 60 and 110 years, respectively. Thus the CFCs are carried by air current slowly into the upper stratosphere (30 to 50 km) where they are decomposed by the strong ultraviolet radiation, producing chlorine atoms. The chlorine atom acts as a catalyst and is extremely destructive to ozone molecules. Even if after all CFCs and other chlorine sources had been cut off in 1996, chlorine levels would continue to rise in the stratosphere for another 20 or 30 years. We now know that there are other ozone layer depleting chemicals. A summary is listed below: Approx. atmos. Approx. Ozone layer depleting chemicals lifetime (years) Contribution CFCs CFC-12 (air conditioning, refrigeration) % CFC-11 (foams, aerosols) 60 26% CFC-113 (solvent) 90 12% Carbon tetrachloride (solvent) 70 8% Methyl chloroform (solvent) 8 5% Halon 1301 (fire extinguishers) 100 4% There are other ozone layer depleting chemicals not listed in the table above. They include methyl bromide, which is used as a soil fumigant, NOx released by supersonic aircrafts and space shuttles, and nitrous oxide (N 2 O) resulting from the use of fertilizers. A chemistry professor at U. C. Irvine, F. Sherwood Rowland, and then postdoctoral student Mario Molina first warned in 1974 that chlorofluorocarbons (CFCs) could destroy the ozone layer in the stratosphere. In 1995, they were awarded, together with German chemist

5 10 5 Paul Crutzen, the Nobel Prize in chemistry for their work toward understanding the formation and destruction of stratospheric ozone. The problem with ozone depletion is that it is potentially life threatening, but like so many other environmental issues, is filled with uncertainties. In 1978, the U.S. banned the use of CFC in spray cans but not much else. The definitive scientific evidence finally came in 1985 when atmospheric scientists of the British Antarctic Survey published their satellite measurements on what is now known as the ozone hole over Antarctica. What we call the hole is figurative. It refers to the region where the stratospheric ozone level has decreased by 30% or more. Still, the ozone hole is as big as the continental U.S. The rate of decline was about 0.5% per year when CFC levels in the stratosphere more than doubled in the 1990s. Uses of CFC have been banned under the Montreal Protocol. In September 1987, 24 industrial nations signed the Montreal Protocol which initially aimed to cut down the use of ozone layer depleting chemicals, which of course, include CFCs. The agreement was amended in 1990 and 1992 to require phasing out these chemicals by the turn of the century. Alarmed by newer reports of ozone layer depletion in the Northern Hemisphere, former President George H.W. Bush announced that the United States would halt production of major ozone layer depleting chemicals by Dec. 31, 1995, rather than by 2000 as set by the Montreal Protocol. The international community agreed and the Montreal Protocol was amended again to adopt the expedited phase-out schedule. EPA scientists had determined that it was technically feasible to phase out production of CFCs and halons by 1997 or earlier, and methyl chloroform and carbon tetrachloride as early as For example, the electronic and aerospace industries had developed new processes to clean their components without using CFCs or other chlorine-containing organic solvents. Scientists with the National Oceanic and Atmospheric Administration reported that ground stations around the world detected for the first time a slight decline (1 to 1.5% from early 1994 to mid-1995) in the amount of ozone-destroying chemicals in the atmosphere. Nonetheless, because of the long lifetimes of CFCs in the atmosphere, the stratospheric ozone hole is not expected to heal until after What's so bad about losing ozone in the stratosphere? Ozone is a paradoxical molecule. At ground level, ozone is a pollutant. Just a bit higher up but still in the troposphere, ozone is a greenhouse gas. In the stratosphere, it plays a crucial role as a sunscreen. With a reduction in stratospheric ozone, an increased amount of harmful ultraviolet radiation will reach the earth. The radiation can increase the incidence of skin cancer. It also suppresses the immune system and causes cataracts. Food crops and aquatic ecosystems are also threatened. Over half of the plants tested have been found to be damaged by ultraviolet light. The growth rate of phytoplankton has been found to decrease when they are exposed to greater amounts of damaging ultraviolet light. The decrease may explain the recent 2% to 4% drop in phytoplankton growth in the Antarctic region. Since phytoplankton is the foundation of the Antarctic food chain, its slower growth will lead to decreases in fish population. The radiation can penetrate underwater as deep as 20 meters under clear conditions, affecting the fisheries directly. There is evidence that during the period when the ozone hole

6 10 6 over Antarctica is most intense, the increased ultraviolet rays are causing DNA damage (lesions) in the eggs and larvae of fish. Phytoplankton that live on the underside of Antarctic sea ice also have suffered reduced reproduction rates in recent years. We may also consider other indirect effects. If increased ultraviolet radiation impedes the growth of green plants, it may worsen the greenhouse effect. In a different aspect, increased ultraviolet radiation may speed up photochemical reactions, exacerbating urban smog. CFCs are so destructive because the chlorine atom that they release participates in a catalytic cycle. Even though CFC molecules are heavier than nitrogen and oxygen molecules (air), the CFCs eventually drift into the upper atmosphere by air current because they are so stable in the lower atmosphere. In the stratosphere, the strong sunlight now has enough energy to dissociate the CFC molecule. The dissociation of CFC releases a chlorine atom. The chlorine atom (Cl) in turn participates in a catalytic cycle: Cl + O 3 > ClO + O 2 O + ClO > Cl + O 2 Chlorine is consumed in the first chemical reaction to form chlorine oxide ClO, but it is regenerated in the second reaction. There is no net loss of the chlorine atom hence the term catalytic. The net effect of these two chemical reactions is the destruction of one ozone molecule. Net: O 3 + O > O 2 + O 2 Roughly, one chlorine atom can destroy 100,000 ozone molecules through this catalytic process. The chlorine atom can be removed, but only by some other much slower chemical reactions. The ozone destruction cycle is amplified by PSC that denitrifies the stratosphere. In addition to the chlorine catalytic cycle, the ozone destruction is amplified by the presence of polar stratospheric clouds (PSC). This mechanism also involves condensed nitric acid HNO 3 in a form that is referred to as nitric acid trihydrate (NAT). Polar stratospheric clouds are formed when water vapor condenses at very cold temperatures near 80 ºC to form tiny micrometer-sized ice particles. Most PSCs are not visible easily, but over the poles, they give beautiful pink and green coloration as a result of forward diffraction of sunlight. Under normal circumstances, the chlorine atom Cl generated by the catalytic cycle can be removed by methane (CH 4 ) to form hydrogen chloride (HCl). Likewise, chlorine oxide ClO can react with nitrogen dioxide (NO 2 ) or nitric acid (HNO 3 ) to form the compound ClONO 2. By and large, a good amount of Cl and ClO can be tied up in the molecules HCl and ClONO 2. These compounds are called inert reservoirs (of chlorine). However, Cl can be released from these insert reservoir molecules in the presence of a solid surface. Up in the stratosphere, the surface is provided by the tiny ice crystals in polar stratospheric clouds which are formed only in very cold temperatures. In the presence of PSC, nitric acid prefers to form nitric acid trihydrate (NAT) on the PSC ice particles. As HNO 3 is removed, the atmosphere is denitrified and loses the gaseous nitric acid that could have been used to form the reservoir species. Worse, ClONO 2 dissociates on

7 10 7 the surface of the PSC, losing the nitrogen to the ice particle to form NAT, while active chlorine is released, free again to destroy ozone molecules. Tiny sulfate particulates (aerosols), spewed forth by volcanic eruptions could create, in effect, temporary polar stratospheric clouds. NASA research satellite pictures showed that the gigantic eruption of Mount Pinatubo in 1991 accelerated the erosion of stratospheric ozone even in tropical regions. Another important chemical process in the stratosphere involves nitrous oxide, N 2 O. Nitrous oxide, produced by the action of soil bacteria on fertilizer and manure, can itself catalyze the break down of ozone. But by and large, nitrous oxide and other chlorinecontaining gases react with each other at lower altitudes (below 40 km), reducing the risk to the ozone layer. Above Antarctica, however, it is so cold that nitrous oxide is frozen out of the atmosphere. The CFCs that are left behind do their damage. What leads to the Antarctic ozone hole? Depletion of stratospheric ozone occurs worldwide. Nevertheless, two phenomena, the polar vortex and the presence of polar stratospheric clouds, contribute to the conspicuous Antarctic ozone hole each year around September and October. Combining what we have learned in the last few sections, here is a quick rundown of the annual events at Antarctica: June July-August September October November Antarctic winter begins. Polar vortex develops. A mass of cold air circulates over Antarctica, isolating it from warmer air beyond the polar region. Temperature drops so low that polar stratospheric clouds are formed. Early spring in Antarctica. Stronger sunlight releases more chlorine atom from CFCs. The polar stratospheric clouds also release chlorine atoms from the inert reservoirs. These events accelerate the destruction of ozone. Temperature rises. The polar stratospheric clouds are gone, but ozone is at the lowest level. Polar vortex breaks up (under normal circumstances), dispersing the ozone depleted atmosphere elsewhere. Ozone layer depletion is linked to climate change. Through the later decades of the 20th Century, warmer temperatures at the outskirt of Antarctica led to the break up of ice shelves. In a perplexing contradiction, the interior of Antarctica cooled. The phenomenon is tied to an anthropogenic input on a natural event. The polar vortex of air current that encircles the Antarctica each winter usually dissipates in the spring. Now, this circumpolar vortex stays intact way into summer time (December and January). By bottling up cold air in the center and restricting warmer air to the exterior, the prolonged vortex would create the perplexing temperature difference. The next question, of course, is why the polar vortex stays longer than usual. The reason lies in the observation that ozone layer depletion also cools the stratosphere by more than 6 ºC because there is much less ozone to absorb the higher energy solar radiation. It is highly probable that the ozone loss-induced cooling propagates downward into the troposphere, thus extending the wintry vortex into the summer months. (Until more recent measurements, scientists had thought that the stratosphere, with a much lighter air mass, is too weak to affect the denser air below it.)