Acid Rain

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Acid rain

"Acid rain" is a popularly used phrase that refers to the deposition of acidifying substances from the atmosphere and the environmental damage that this causes. Acid rain became a prominent issue around 1970, and since then research has demonstrated that the deposition of atmospheric chemicals is causing widespread acidification of lakes and streams, and possibly soil . The resulting biological effects include the extirpation (or local extinction ) of many populations of fish . Scientific understanding of the causes and consequences of acid rain, in conjunction with lobbying of government by environmental organizations, has resulted in large reductions in the atmospheric emissions of pollutants in North America and parts of Europe . If these reductions prove to be large enough, acid rain will be less of an environmental problem in those regions.

Atmospheric deposition

Strictly speaking, the term "acid rain" should only refer to rainfall, or so-called wet precipitation . However, the proper meaning of acid rain is "the deposition of acidifying substances from the atmosphere." This is because acidification is not just caused by acidic rain, but also by chemicals in snow and fog , and by inputs of gases and particulates when precipitation is not occurring.

Of the many chemicals that are deposited from the atmosphere, the most important in terms of causing acidity in soil and surface waters (such as lakes and streams) are: (1) dilute solutions of sulfuric and nitric acids (H2SO4 and HNO3, respectively) deposited as acidic rain or snow, (2) the gases sulfur dioxide (SO2) and oxides of nitrogen (NO and NO2, together called NOx), and (3) tiny particulates, such as ammonium sulfate ([NH4]2SO4) and ammonium nitrate (NH4NO3).

The depositions of these gases and particulates primarily occur when it is not raining or snowing. This type of atmospheric input is known as "dry deposition." Large regions of Europe and North America are exposed to these acidifying depositions. However, only certain types of ecosystems are vulnerable to becoming acidified by these atmospheric inputs. These usually have a thin cover of soil that contains little calcium , and sits upon a bedrock of hard minerals such as granite or quartz. There is convincing evidence that atmospheric depositions have caused an acidification of freshwater ecosystems in such areas. Many lakes, streams, and rivers have become acidic, resulting in declining or locally extirpated populations of some plants and animals. However, there is not yet conclusive evidence that terrestrial ecosystems have been degraded by acidic deposition (except for cases of severe pollution by toxic SO2).

Chemistry of precipitation

The acidity of an aqueous solution is measured as its concentration of hydrogen ions (H+). The pH scale expresses this concentration in logarithmic units to the base 10, ranging from very acidic solutions of pH 0, through the neutral value of pH 7, to very alkaline (or basic) solutions of pH 14. It is important to recognize that a one-unit difference in pH (for example, from pH 3 to pH 4) implies a 10-fold difference in the concentration of hydrogen ions. The pHs of some common solutions include: lemon juice, pH 2; table vinegar, pH 3; milk, pH 6.6; milk of magnesia, pH 10.5.

As just noted, an acidic solution, strictly speaking, has a pH less than 7.0. However, in environmental science the operational definition of acidic precipitation is a pH less than 5.65. This is the pH associated with the weak solution of carbonic acid (H2CO3) that forms when water droplets in clouds are in chemical equilibrium with carbon dioxide (CO2), an atmospheric gas with a concentration of about 360 ppm (parts per million; this is a unit of concentration).

Water in precipitation contains a mixture of positively charged ions (or cations) and negatively charged ions (or anions). The most abundant cations are usually hydrogen (H+), ammonium (NH4+), calcium (Ca2+), magnesium (Mg2+), and sodium (Na+), while the major anions are sulfate (SO42-), chloride (Cl-), and nitrate (NO3-). The principle of conservation of electrochemical neutrality of aqueous solutions states that the total number of cation charges must equal that of anions, so the net electrical charge is zero . Following from this principle, the quantity of H+ in an aqueous solution is related to the difference in concentration of the sum of all anions, and the sum of all cations other than H+.

Data for the chemistry of precipitation in a region experiencing severe acid rain are available from Hubbard Brook, New Hampshire, where one of the world's best long-term studies of this phenomenon has been undertaken. The average pH of precipitation at Hubbard Brook is 4.2, and H+ accounts for 71% of the total amount of cations, and SO42- and NO3- for 87% of the anions. Therefore, most of the acidity of precipitation at Hubbard Brook occurs as dilute sulfuric and nitric acids. The SO42- is believed to originate from SO2 emitted from power plants and industries, and oxidized by photochemical reactions in the atmosphere to SO42-. The NO3- originates with emissions of NOx (i.e., NO and NO2) gases from these sources and automobiles. Not surprisingly, air masses that pass over the large emission sources of Boston and New York produce storms with the highest concentrations of H+, SO42-, and NO3- at Hubbard Brook.

Regions differ greatly in their precipitation chemistry. This can be demonstrated using data for precipitation chemistry monitored during a study in eastern Canada. The village of Dorset in southern Ontario is close to large sources of emission of SO2 and NOx. On average, the precipitation at Dorset is highly acidic at pH 4.1, and the large concentrations of SO42- and NO3- suggest that the acidity is caused by dilute sulfuric and nitric acids. In comparison, the Experimental Lakes Area (ELA) is in a remote landscape in northwestern Ontario that is infrequently affected by polluted air masses. The ELA site has a less acidic precipitation (average pH 4.7) and smaller concentrations of SO42- and NO3- than at Dorset. Another site near the Atlantic Ocean in Nova Scotia receives air masses that pass over large sources of emissions in New England and southeastern Canada. However, by the time Nova Scotia is reached much of the acidic SO42- and NO3- have been removed by prior rain-out, and the precipitation is only moderately acidic (pH 4.6). Also, because Nova Scotia is influenced by the ocean, its precipitation chemistry is characterized by high concentrations of Na+ and Cl-. Finally, Lethbridge in southern Alberta is in a prairie landscape, and its precipitation is not acidic (average pH 6.0) because of the influence of calcium-rich, acid-neutralizing dusts blown into the atmosphere from agricultural fields.

In some places, fog moisture can be especially acidic. For example, fogwater at coastal locations in New England can be as acidic as pH 3.0-3.5. At high-elevation locations where fog is frequent there can be large depositions of cloudwater and acidity. At a site in New Hampshire where fog occurs 40% of the time, cloudwater deposition to a conifer forest is equivalent to 47% of the water input by rain and snow, and because of its large concentrations of some chemicals, fog deposition accounted for 62% of the total inputs of H+, and 81% of those of SO42- and NO3-.

Spatial patterns of acidic precipitation

Large regions are affected by acidic precipitation in North America, Europe, and elsewhere. A relatively small region of eastern North America is known to have experienced acidic precipitation before 1955, but this has since expanded so that most of the eastern United States and southeastern Canada is now affected.

Interestingly, the acidity of precipitation is not usually greater close to large point-sources of emission of important gaseous precursors of acidity, such as smelters or power plants that emit SO2 and NOx. This observation emphasizes the fact that acid rain is a regional phenomenon, and not a local one. For instance, the acidity of precipitation is not appreciably influenced by distance from the world's largest point-source of SO2 emissions, a smelter in Sudbury, Ontario. Furthermore, when that smelter was temporarily shut down by a labor dispute, the precipitation averaged pH 4.49, not significantly different from the pH 4.52 when there were large emissions of SO2.

Dry deposition of acidifying substances

Dry deposition occurs in the intervals of time between precipitation events. Dry deposition includes inputs of tiny particulates from the atmosphere, as well as the uptake of gaseous SO2 and NOx by plants, soil, and water. Unlike wet deposition, the rates of dry deposition can be much larger close to point-sources of emission, compared with further away.

Once they are dry deposited, certain chemicals can generate important quantities of acidity when they are chemically transformed in the receiving ecosystem . For example, SO2 gas can dissolve into the water of lakes or streams, or it can be absorbed by the foliage of plants. This dry-deposited SO2 is then oxidized to SO42-, which is electrochemically balanced by H+, so that acidity results. Dry-deposited NOx gas can similarly be oxidized to NO3- and also balanced by H+.

In relatively polluted environments close to emissions sources, the total input of acidifying substances (i.e., wet + dry depositions) is dominated by the dry deposition of acidic substances and their acid-forming precursors. The dry deposition is mostly associated with gaseous SO2 and NOx, because wet deposition is little influenced by distance from sources of emission.

For example, within a 25 mi (40 km) radius of the large smelter at Sudbury, about 55% of the total input of sulfur from the atmosphere is due to dry deposition, especially SO2. However, less than 1% of the SO2emission from the smelter is deposited in that area, because the tall smokestack is so effective at widely dispersing the emissions.

Because they have such a large surface area of foliage and bark , forests are especially effective at absorbing atmospheric gases and particles. Consequently, dry inputs accounted for about 33% of the total sulfur deposition to a hardwood forest in New Hampshire, 56-63% of the inputs of S and N to a hardwood forest in Tennessee, and 55% of their inputs to a conifer forest in Sweden.

Chemical changes in the forest canopy

In any forest, leaves and bark are usually the first surfaces encountered by precipitation. Most rainwater penetrates the foliar canopy and then reaches the forest floor as so-called throughfall, while a smaller amount runs down tree trunks as stemflow. Throughfall and stemflow have a different chemistry than the original precipitation. Because potassium is easily leached out of leaves, its concentration is especially changed. In a study of several types of forest in Nova Scotia, the concentration of potassium (K+) was about 10 times larger in throughfall and stemflow than in rain, while calcium (Ca2+) and magnesium (Mg2+) were three to four times more concentrated. There was less of a change in the concentration of H+; the rainwater pH was 4.4, but in throughfall and stemflow of hardwood stands pH averaged 4.7, and it was 4.4-4.5 in conifer stands. The decreases in acidity were associated with ion-exchange reactions occurring on foliage and bark surfaces, in which H+ is removed from solution in exchange for Ca2+, Mg2+, and K+. Overall, the "consumption" of hydrogen ions accounted for 42-66% of the input of H+ by precipitation to these forests. Similarly, H+ consumption by the tree canopy was 91% in a hardwood forest at Hubbard Brook, New Hampshire, 21-80% among seven stands in New Brunswick, and 14-43% in stands in upstate New York.

In areas polluted by SO2 there can be large increases in the sulfate concentration of throughfall and stemflow, compared with ambient precipitation. This is caused by the washoff of SO2 and SO4 that had been previously dry-deposited to the canopy. At Hubbard Brook this SO4 enhancement is about four times larger than ambient precipitation, while in central Germany it is about two to three times greater. These are both regions with relatively large concentrations of particulate SO4 and gaseous SO2 in the atmosphere.

Chemical changes in soil

Once precipitation reaches the forest floor, it percolates into the soil. Important chemical changes take place as: (1) microbes and plants selectively absorb, release, and metabolize chemicals; (2) ions are exchanged at the surfaces of particles of clay and organic matter ; (3) minerals are made soluble by so-called acid-weathering reactions; and (4) secondary minerals such as certain clays and metal oxides are formed through chemical precipitation of soluble ions of aluminum , iron , and other metals. These various chemical changes can contribute to: soil acidification, the leaching of important chemicals such as calcium and magnesium, and the mobilization of toxic ions of aluminum, especially Al3+. These are all natural, closely linked processes, occurring wherever there is well-established vegetation, and where water inputs by precipitation are greater than evapotranspiration (i.e., evaporation from vegetation and non-living surfaces). A potential influence of acid rain is to increase the rates of some of these processes, such as the leaching of toxic H+ and Al3+ to lakes and other surface waters.

Some of these effects have been examined by experiments in which simulated "rainwater" of various pHs was added to soil contained in plastic tubes. These experiments have shown that very acidic solutions can cause: (1) an acidification of the soil; (2) increased leaching of the so-called "basic cations" Ca, Mg, and K, resulting in nutrient loss, decreased base saturation of cation exchange capacity, and increased vulnerability of soil to acidification; (3) increased solubilization of toxic ions of metals such as aluminum, iron, manganese, lead , and zinc; and (4) saturation of the ability of soil to absorb sulfate, after which sulfate leaches at about the rate of input. The leaching of sulfate has a secondary influence on soil acidification if it is accompanied by the loss of base cations, and it can cause acidifying and toxic effects in surface waters if accompanied by Al3+ and H+.

Soil acidification can occur naturally. This fact can be illustrated by studies of ecological succession on newly exposed parent materials of soil. At Glacier Bay, Alaska, the melting of glaciers exposes a mineral substrate with a pH of about 8.0, with up to 7-10% carbonate minerals. As this material is colonized and modified by vegetation and climate, its acidity increases, reaching about pH 4.8 after 70 years when a conifer forest has established. Accompanying this acidification is a reduction of carbonates to less than 1%, caused by leaching and uptake by plants.

Several studies have attempted to determine whether naturally occurring soil acidification has been intensified as a result of acid rain and associated atmospheric depositions. So far, there is no conclusive evidence that this has occurred on a wide scale. It appears that soil acidification is a potential, longer-term risk associated with acid rain.

Chemistry of surface waters

Compared with the water of precipitation, that of lakes, ponds, streams, and rivers is relatively concentrated in ions, especially in calcium, magnesium, potassium, sodium, sulfate, and chloride. These chemicals have been mobilized from the terrestrial part of the watersheds of the surface waters. In addition, some surface waters are brown-colored because of their high concentrations of dissolved organic compounds, usually leached out of nearby bogs. Brown-water lakes are often naturally acidic, with a pH of about 4 to 5.

Seasonal variations in the chemistry of surface waters are important. Where a snowpack accumulates, meltwater in the springtime can be quite acidic. This happens because soils are frozen and/or saturated during snowmelt, so there is little possibility to neutralize the acidity of meltwater. So-called "acid shock" events in streams have been linked to the first meltwaters of the snowpack, which are generally more acidic than later fractions.

A widespread acidification of weakly-buffered waters has affected the northeastern United States, eastern Canada, Scandinavia, and elsewhere. In 1941, for example, the average pH of 21 lakes in central Norway was 7.5, but only 5.4-6.3 in the 1970s. Before 1950 the average pH of 14 Swedish water bodies was 6.6, but 5.5 in 1971. In New York's Adirondack Mountains, 4% of 320 lakes had pH less than 5 in the 1930s, compared with 51% of 217 lakes in that area in 1975 (90% were also devoid of fish). The Environmental Protection Agency sampled a large number of lakes and streams in the United States in the early 1990s. Out of 10,400 lakes, 11% were acidic, mostly in the eastern United States. Atmospheric deposition was attributed as the cause of acidification of 75% of the lakes, while 3% had been affected by acidic drainage from coal mines, and 22% by organic acids from bogs. Of the 4,670 streams considered acidic, 47% had been acidified by atmospheric deposition, 26% by acid-mine drainage, and 27% by bogs.

Surface waters that are vulnerable to acidification generally have a small acid-neutralizing capacity. Usually, H+ is absorbed until a buffering threshold is exceeded, and there is then a rapid decrease in pH until another buffering system comes into play. Within the pH range of 6 to 8, bicarbonate alkalinity is the natural buffering system that can be depleted by acidic deposition. The amount of bicarbonate in water is determined by geochemical factors, especially the presence of mineral carbonates such as calcite (CaCO3) or dolomite (Ca,MgCO3) in the soil, bedrock, or aquatic sediment of the watershed . Small pockets of these minerals are sufficient to supply enough acid-neutralizing capacity to prevent acidification, even in regions where acid rain is severe. In contrast, where bedrock, soil, and sediment are composed of hard minerals such as granite and quartz, the acid-neutralizing capacity is small and acidification can occur readily. Vulnerable watersheds have little alkalinity and are subject to large depositions of acidifying substances; these are especially common in glaciated regions of eastern North America and Scandinavia, and at high altitude in more southern mountains (such as the Appalachians) where crustal granite has been exposed by erosion .

High-altitude, headwater lakes and streams are often at risk because they usually have a small watershed. Because there is little opportunity for rainwater to interact with the thin soil and bedrock typical of headwater systems, little of the acidity of precipitation is neutralized before it reaches surface water.

In overview, the acidification of freshwaters can be described as a titration of a dilute bicarbonate solution with sulfuric and nitric acids derived from atmospheric deposition. In waters with little alkalinity, and where the watershed provides large fluxes of sulfate accompanied by hydrogen and aluminum ions, the waterbody is vulnerable to acidification.

Effects of acidification on terrestrial plants

Few studies have demonstrated injury to terrestrial plants caused by an exposure to ambient acid rain. Although many experiments have demonstrated injury to plants after treatment with artificial "acid rain" solutions, the toxic thresholds are usually at substantially more acidic pHs than normally occur in nature.

For example, some Norwegian experiments involved the treating of young forests with simulated acid rain. Lodgepole pine watered for three years grew 15-20% more quickly at pHs 4 and 3, compared with a "control" treatment of pH 5.6-6.1. The height growth of spruce was not affected over the pH range 5.6 to 2.5, while Scotch pine was stimulated by up to 15% at pHs of 2.5 to 3.0, compared with pH 5.6-6.1. Birch trees were also stimulated by the acid treatments. However, the feather mosses that dominated the ground vegetation were negatively affected by acid treatments.

Because laboratory experiments can be well controlled, they are useful for the determination of dose-response effects of acidic solutions on plants. In general, growth reductions are not observed unless treatment pHs are more acidic than about 3.0, and some species are stimulated by more acidic pHs than this. In one experiment, the growth of white pine seedlings was greater after treatment at pHs of 2.3 to 4.0 than at pH 5.6. In another experiment, seedlings of 11 tree species were treated over the pH range 2.6 to 5.6. Injuries to foliage occurred at pH 2.6, but only after a week of treatment with this very acidic pH.

Overall, it appears that trees and other vascular plants are rather tolerant of acidic rain, and they may not be at risk of suffering direct, short-term injury from ambient acidic precipitation. It remains possible, however, that even in the absence of obvious injuries, stresses associated with acid rain could decrease plant growth. Because acid rain is regional in character, these yield decreases could occur over large areas, and this would have important economic implications. This potential problem is most relevant to forests and other natural vegetation. This is because agricultural land is regularly treated with liming agents to reduce soil acidity, and because acid production by cropping and fertilization is much larger than that caused by atmospheric depositions.

Studies in western Europe and eastern North America have examined the possible effects of acid rain on forest productivity. Recent decreases in productivity have been shown for various tree species and in various areas. However, progressive decreases in productivity are natural as the canopy closes and competition intensifies in developing forests. So far, research has not separated clear effects of regional acid rain from those caused by ecological succession, insect defoliation, or climate change.

Effects of acidification on freshwater organisms

The community of microscopic algae (or phytoplankton ) of lakes is quite diverse in species. Non-acidic, oligotrophic (i.e., unproductive) lakes in a temperate climate are usually dominated by golden-brown algae and diatoms , while acidic lakes are typically dominated by dinoflagellates, cryptomonads, and green algae.

An important experiment was performed in a remote lake in Ontario, in which sulfuric acid was added to slowly acidify the entire lake, ultimately to about pH 5.0 from the original pH of 6.5. During this whole-lake acidification, the phytoplankton community changed from an initial domination by golden-brown algae to dominance by green algae. There was no change in the total number of species, but there was a small increase in algal biomass after acidification because of an increased clarity of the water.

In some acidified lakes the abundance of larger plants (called macrophytes) has decreased, sometimes accompanied by increased abundance of a moss known as Sphagnum. In itself, proliferation of Sphagnum can cause acidification, because these plants efficiently remove cations from the water in exchange for H+, and their mats interfere with acid neutralizing processes in the sediment.

Zooplankton are small crustaceans living in the water column of lakes. These animals can be affected by acidification through: (1) the toxicity of H+ and associated metals ions, especially Al3+; (2) changes in their phytoplankton food; and (3) changes in predation, especially if plankton-eating fish become extirpated by acidification. Surveys have demonstrated that some zooplankton species are sensitive to acidity, while others are more tolerant. In general, higher-pH lakes are richer in zooplankton species. For example, a survey of lakes in Ontario found 9-16 species with three to four dominants at pH greater than pH 5, but only 1-7 species with one to two dominants at more acidic pHs.

In the whole-lake experiment mentioned previously, the abundance of zooplankton increased by 66-93% after acidification, a change attributed to an increase in algal biomass. Although there was little change in dominant species, some less common species were extirpated.

Fish are the best-known victims of acidification. Loss of populations of trout, salmon , and other species have occurred in many acidified freshwaters. A survey of 700 Norwegian lakes, for example, found that brown trout were absent from 40% of the water bodies and sparse in another 40%, even though almost all of the lakes had supported healthy fish populations prior to the 1950s. Surveys during the 1930s in the Adirondack Mountains of New York found brook trout in 82% of the lakes. However, in the 1970s fish did not occur in 43% of 215 lakes in the same area, including 26 definite extirpations of brook trout in re-surveyed lakes. This dramatic change paralleled the known acidification of these lakes. Other studies documented the loss of fish populations from lakes in the Killarney region of Ontario, where there are known extirpations of lake trout in 17 lakes, while smallmouth bass have disappeared from 12 lakes, largemouth bass and walleye from four, and yellow perch and rock bass from two.

Many studies have been made of the physiological effects of acidification on fish. Younger life-history stages are generally more sensitive than adults, and most losses of fish populations can be attributed to reproductive failure, rather than mortality of adults (although adults have sometimes been killed by acid-shock episodes in the springtime).

There are large increases in concentration of certain toxic metals in acidic waters, most notably ions of aluminum. In many acidic waters aluminium ions can be sufficient to kill fish, regardless of any direct effect of H+. In general, survival and growth of larvae and older stages of fish are reduced if dissolved aluminium concentrations are larger than 0.1 ppm, an exposure regularly exceeded in acidic waters. The most toxic ions of aluminium are Al3+ and AlOH2+.

Although direct effects of acidification on aquatic birds have not been demonstrated, changes in their habitat could indirectly affect their populations. Losses of fish populations would be detrimental to fish-eating waterbirds such as loons , mergansers, and osprey. In contrast, an increased abundance of aquatic insects and zooplankton, resulting from decreased predation by fish, could be beneficial to diving ducks such as common goldeneye and hooded merganser, and to dabbling ducks such as the mallard and black duck.

Reclamation of acidified water bodies

Fishery biologists especially are interested in liming acidic lakes to create habitat for sportfish. Usually, acidic waters are treated by adding limestone (CaCO3) or lime (Ca[OH]2), a process analogous to a whole-lake titration to raise pH. In some parts of Scandinavia liming has been used extensively to mitigate the biological damages of acidification. By 1988 about 5,000 water bodies had been limed in Sweden, mostly with limestone, along with another several hundred lakes in southern Norway. In the early 1980s there was a program to lime 800 acidic lakes in the Adirondack region of New York.

Although liming rapidly decreases the acidity of a lake, the water later re-acidifies at a rate determined by size of the drainage basin , the rate of flushing of the lake, and continued atmospheric inputs. Therefore, small headwater lakes have to be re-limed more frequently. In addition, liming initially stresses the acid-adapted biota of the lake, causing changes in species dominance until a new, steady-state ecosystem is achieved. It is important to recognize that liming is a temporary management strategy, and not a long-term solution to acidification.

Avoiding acid rain

Neutralization of acidic ecosystems treats the symptoms, but not the sources of acidification. Clearly, large reductions in emissions of the acid-forming gases SO2 and NOx are the ultimate solution to this widespread environmental problem. However, there is controversy over the amount that the emissions must be reduced in order to alleviate acidic deposition, and about how to pursue the reduction of emissions. For example, should large point sources such as power plants and smelters be targeted, with less attention paid to smaller sources such as automobiles and residential furnaces? Not surprisingly, industries and regions that are copious emitters of these gases lobby against emission controls, for which they argue the scientific justification is not yet adequate.

In spite of many uncertainties about the causes and magnitudes of the damage associated with acid rain and related atmospheric depositions, it is intuitively clear that what goes up (that is, the acid-precursor gases) must come down (as acidifying depositions). This common-sense notion is supported by a great deal of scientific evidence, and because of public awareness and concerns about acid rain in many countries, politicians have began to act effectively. Emissions of sulfur dioxide and oxides of nitrogen are being reduced, especially in western Europe and North America. For example, in 1992 the governments of the United States and Canada signed an air-quality agreement aimed at reducing acidifying depositions in both countries. This agreement calls for large expenditures by government and industry to achieve substantial reductions in the emissions of air pollutants during the 1990s. Eventually, these actions should improve environmental conditions related to damage caused by acid rain.

However, so far the actions to reduce emissions of the precursor gases of acidifying deposition have only been vigorous in western Europe and North America. Actions are also needed in other, less wealthy regions where the political focus is on industrial growth, and not on control of air pollution and other environmental damages that are used to subsidize that growth. In the coming years, much more attention will have to be paid to acid rain and other pollution problems in eastern Europe and the former USSR, China, India, southeast Asia , Mexico, and other so-called "developing" nations. Emissions of important air pollutants are rampant in these places, and are increasing rapidly.

Resources

books

Edmonds, A. Acid Rain. Sussex, England: Copper Beech Books, Ltd., 1997.

Ellerman, Danny. Markets for Clean Air: The U. S. Acid RainProgram. Cambridge: Cambridge University Press, 2000.

Freedman, B. Environmental Ecology. 2nd ed. San Diego: Academic Press, 1995.

Hancock P. L. and Skinner B. J., eds. The Oxford Companion to the Earth. Oxford: Oxford University Press, 2000.

periodicals

Anonymous. National Acid Precipitation Assessment Program Integrated Assessment Report. Washington, DC: Superintendent of Documents, U.S. Government Printing Office, 1989.

Brimblecombe, P. "Acid Rain 2000." Water, Air, and Soil Pollution 130, 1-4 (2001): 25-30.

Galloway, James N. "Acidification of the World: Natural and Anthropogenic." Water, Air, and Soil Pollution 130, no. 1-4 (2001): 17-24.

Krajick, K. "Acid Rain: Long-term Data Show Lingering Effects from Acid Rain." Science 292, no. 5515 (2001): 195-196.

Milius, S. "Red Snow, Green Snow." Science News no. 157 (May 2000): 328-333.

other

The United Nations. "The Conference and Kyoto Protocol," homepage [cited March 2003]. <http://unfccc.int/resource/convkp.html>.

United Stated Geological Survey. "What is Acid Rain?" [cited March 2003]. <http://pubs.usgs.gov/gip/acidrain/2.html>.

Bill Freedman

KEY TERMS

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Acid mine drainage: —Surface water or groundwater that has been acidified by the oxidation of pyrite and other reduced-sulfur minerals that occur in coal and metal mines and their wastes.
Acid shock: —A short-term event of great acidity. This phenomenon regularly occurs in freshwater systems that receive intense pulses of acidic water when an accumulated snowpack melts rapidly in the spring.
Acidic rain (acidic precipitation): —(1) Rain, snow, sleet or fog water having a pH less than 5.65. (2) The deposition of acidifying substances from the atmosphere during a precipitation event.
Acidification: —An increase over time in the content of acidity in a system, accompanied by a decrease in acid-neutralizing capacity.
Acidifying substance: —Any substance that causes acidification. The substance may have an acidic character and therefore act directly, or it may initially be non-acidic but generate acidity as a result of its chemical transformation, as happens when ammonium is nitrified to nitrate, and when sulfides are oxidized to sulfate.
Acidity: —The ability of a solution to neutralize an input of hydroxide ion (OH—). Acidity is usually measured as the concentration of hydrogen ion (H+), in logarithmic pH units (see also pH). Strictly speaking, an acidic solution has a pH less than 7.0.
Acidophilous: —Refers to organisms that only occur in acidic habitats, and are tolerant of the chemical stresses of acidity.
Conservation of electrochemical neutrality: —Refers to an aqueous solution, in which the number of cation equivalents equals the number of anion equivalents, so that the solution does not have a net electrical charge.
Equivalent: —Abbreviation for mole-equivalent, and calculated as the molecular or atomic weight multiplied times the number of charges of the ion. Equivalent units are necessary for a charge-balance calculation, related to the conservation of electrochemical neutrality (above).
Leaching: —The movement of dissolved chemicals with water percolating through soil.
pH: —The negative logarithm to the base 10 of the aqueous concentration of hydrogen ions in units of moles per liter. An acidic solution has pH less than 7, while an alkaline solution has pH greater than 7. Note that a one-unit difference in pH implies a 10-fold difference in the concentration of hydrogen ions.