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8 . AC ID PREC I PITATION The deposition of acid from the atmosphere was recognized in President Carter's second environmental message to the U.S. Congress on August 2, 1979, as "one of the most serious global pollution problems associated with fossil fuel combustion," rivaled only by the buildup of carbon dioxide in the atmosphere. The topic, generally known as "acid rain," has been of much international concern because acids are deposited far from the sources of their precursors. Acid precipitation was the subject of a major international symposium in 1975, when research on many aspects of the topic was in its infancy. Since that time, many studies have been completed, including an enormous amount of work in Scandinavia. Recent symposia (Drablos and Tollan 1980, Hutchinson and Havas 1980, Shriner et al. 1980) treat many aspects of the problem in much more detail than previously, and some long-term studies of effects allow a much more conclusive summary of the problem. We can thus learn much from a case study of acid precipitation about the problems likely to be encountered with other atmospheric pollutants. CAUSES OF ACID PREC IPITATION It is thought that the pH of "pure" rain is controlled by the weak acid, carbonic acid (H2CO3), resulting from atmospheric CO2 in solution. The resulting pH would be near 5.6. When alkaline dust and ocean sea spray are taken into account, the pH of precipitation must be higher than pH 5.6, probably around pH 7.0 (Oden 1976~. But this theory is no longer testable, because global pollution of the atmosphere with sulfur compounds from fossil fuel combustion and the smelting of metalliferous sulfide ores has existed for decades, if not centuries, as shown by analyses of polar icecaps (Figure 8.1~. It seems likely that under unpolluted conditions, small releases of sulfur and nitrogen oxides from volcanic acitivity and microbial metabolism to the biosphere would cause slightly lower natural pH levels than the theoretical level in geologically unbuffered areas, and that wind-carried calcareous dust would cause pH levels higher 140

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142 than 5.6 in areas where calcareous rocks and soils are plentiful. It is now clear that precipitation is far more acid than theoretically pure rain in regions to which prevailing winds carry oxides of sulfur and nitrogen from heavily industrialized areas (Oden 1968, Dovland and Semb 19801. Direct cause-and-effect linkages between sources of acid and effects on ecosystems will not be possible in the foreseeable future, owing to the remoteness of sources and the complexity of the interaction among emissions from different sources, atmospheric transport, chemical transformations, and specific orographic and geological settings (Table 8.1~. But the increased emission of sulfur and nitrogen compounds from anthropogenic sources is the only plausible explanation for acid deposition. The oxides of these elements~appear to be oxidized further in the atmosphere to form the strong acids, sulfuric acid (H2SO4) and nitric acid (HNO3), which contaminate wet precipitation and atmospheric aerosols (Figure 8.2~. The dry deposition of ammonium sulfate aerosols, which are then converted to sulfuric acid in ecosystems through biological processes, may also contribute substantially to the problem (Oden 1976~. While scientists are in general agreement that industrial emissions of sulfur and nitrogen oxides have caused the contamination of precipitation with strong mineral acids, the timing of the increase and the present rate of increase in acidity are matters of some dispute. Several authors have claimed that the acidity of precipitation has increased rapidly in the past few decades (for example, Cogbill and Likens 1974, Oden and Ahl 1970, Dickson 1975; see Figure 8.3~. This evidence has been disputed by others, who claim that the apparent increase in acidity of precipitation is due to methodological changes (Hansen et al. 1981~. A number of changes in the emissions of acid precursors have taken place over the past few decades, which may influence the acidity of precipitation. While SO2 emissions have not changed greatly for several decades, owing to a switch from coal to other fuels and to increased control of gaseous sulfur emissions, there was unquestionably a great increase in anthropogenic sulfur emissions in the 20th century, causing increased sulfate deposition in remote regions (see Figure 8.1~. Despite the relative constancy of annual SO2 emissions during the past century, three technological changes may cause the emissions to produce acid precipitation more efficiently. First, the height at which gases are injected into the atmosphere has increased nearly threefold (Fig. 8.4), causing SO2 to be transported farther and to remain in the atmosphere longer, increasing the probability of oxidation to sulfuric acid. Second, recent controls of particulate emissions have reduced the amount of alkaline fly ash discharged from smokestacks, and it is conceivable that in the past such material partially neutralized acid emissions. Third, there has been a gradual change from seasonal to year-round emission. Less coal is used for space heating and more coal is used for generating electricity. Also the demand for electricity during the summer months

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143 TABLE 8.1 Factors affecting the vulnerability of an ecosystem to acid rain A. Anthropogenic 1. Spatial and temporal patterns of urban/industrial development 2. Kinds and amounts of energy resources in use 3. Controls on atmospheric emissions 4. Degree of agricultural activity (cultivation, liming, fertilization) B. Geologic 1. Nature of bedrock, as regards both basic minerals and acid-soluble toxic metals 2. Patterns of glaciation 3. Depth, texture, mineralogy, and organic content of soil C. Climatic Amount of precipitation Atmospheric humidity, as it affects gas absorption and particle collision 3. Direction and speed of winds and air-mass movements 4. Temperature, especially as it affects the proportions of rain and snow, and rates of chemical reaction in the atmosphere Ratio of precipitation to evaporation, as it affects leaching and the residence time of water in lakes D. Topographic 1. 2. Altitude, as it influences soil depth, precipitation, etc. Order of streams and lakes in the hydrologic network 3. Lake depth and ratio of watershed area to lake area, controlling residence time of water E. Biotic 1. Height, type, and duration of leaf canopy 2. Magnitude of transpiration 3. Sensitivity of critical species, including the microbes mediating biogeochemical cycles F. Natural, episodic 2. 4. 1. Volcanoes, producing locally acid rain Fires in deposits of fossil fuel such as coal or lignite Forest fires, entraining alkaline particulates into the atmosphere Dust storms, entraining alkaline soil particles into the atmosphere

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144 IV C) o/ I 2' ._ .= 'a J N O O C/' C/) O ~ o ~ .m o 4— Cot ._ 1~ 1° \ Cal \ O \ c o CD —\ ._ X o o Cat ._ C) Cal Cal o Cal Cal ._ o o ._ en o . ~ ~ \ O ~ O en _ O an \ - ~n~ _~ 4) Go -

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145 6-.0 Q 5.8 5.4 pH 5.0 4.6 4.2 5.8 5.4 pH 5.0 4.6 4.2 5.8 5.4 pH 5.0 5.0 4.2 I I 1 1 1955 1960 1965 1970 YEAR (a) — Roba~ksdalen I_. Kise 5.8 5.4 5 . _ Flahult ~ I (Jonkoping) 42 _ A_ 5.8 5.4 5.0 4.6 4 2 5.4 5.0 4.6 - 1 1'.\ ~ 4.6 4.2 1 1 ~ —~ 1 1955 1960 1965 1970 YEAR (b) FIGURE 8.3 The pH levels of precipitation in Scandinavia, 1955-1975. SOURCES: (a) Dickson (1975); (b) Oden and Ahl (1970). /li ~ Pionninge ~ (Halmstad) - T~ ~ .~- Smedby it_ . I I I 1 1955 1960 1965 1970 YEAR

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1~76 1 400 - ~ 1000 c o ID o ~3 al 600 - - ._ I 200 , —— ,, Tallest , ~ , , ~ , \ , \. ~ ~ ~ ~ _ l I; \, ~ stack - Average stack height ,........... 1956 1960 1964 1968 1972 1976 FIGURE 8.4 Average stack height and tallest stack reported among power plants burning fossil fuels (bituminous coal, lignite, oil) included in biannual design surveys of new power plants, 1956-1978. SOURCE: Patrick et al. (1981). Reprinted with per- mission from Science 211:446448. Copyright O 1981 by the American Association for the Advance- ment of Science.

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147 has grown with the increased use of air conditioning. The high temperatures and humidities in summer may result in more efficient oxidation of SO2 emissions to sulfuric acid. In addition to the technological changes in SO2 emission, there has been an increase in emission of nitrogen oxides for the past few decades (see Table 4.2~. Nitrogen oxides are emitted from a wide variety of sources, with some injected high into the atmosphere while others are ejected and dispersed at ground level through motor vehicle use. In the absence of control technology for nitrogen oxides, their emissions will exceed emission of sulfur oxides by the turn of the century. We stress that emission of nitrogen and sulfur oxides and the consequent acid precipitation are broad regional rather than global problems. When natural and anthropogenic emissions are compared on a regional basis, it is clear that man's acitivities completely overwhelm natural sources of SO2 and NOX, (see Table 4.1), even though the magnitude of anthropogenic emissions of these oxides may seem unimportant when compared with natural emissions on a global scale. The observed recent increases in lake acidity could have resulted either from a rapid increase in acid precipitation in recent time or from long-term, constant acid precipitation over several decades duration. It is difficult to differentiate between these two possible patterns because of the nature of the bicarbonate buffering curve. As a solution of bicarbonate--such as a lake--is titrated by a constant addition of strong acid there is little resulting change in pH until 80 to 90 percent of the bicarbonate has been consumed according to the reaction: H+ + HCO3 -, H2CO3 ~ H2O + CO2 Once the bicarbonate has been converted into carbon dioxide and lost to the atmosphere, the acid (hydrogen ions) accumulates and the pH decreases rapidly (Figure 8.5~. The relatively sudden drop in pH to an acid condition is the same regardless whether the titration (the addition of acid) occurred at a fast or slow rate. The lake will become acidic so long as the rate at which acid is added to the lake exceeds the rate at which geochemical weathering processes replace the bicarbonate. The theory that the acidification observed in poorly buffered fresh waters was due to changing land-use patterns (Rosenqvist 1978a,b) has now been discounted as an explanation for the widespread effects observed, particularly in remote areas. Detailed study over several years of watersheds in Norway, some with changing land-use patterns and some without, has shown that, on the average, both are acidified at equal rates (Drablos and Sevaldrud 1980, Drablos et al. 1980~. Moreoever, studies of lakes in North America in areas where land-use patterns have never changed have also shown substantial increases in hydrogen ion or losses in buffering capacity (Dillon et al. 1978, Watt et al. 1979~. Along with the hydrogen ions supplied by transformation of atmospheric sulfur dioxide and oxides of nitrogen to their soluble,

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148 Bir:~rhc~n~t`? Transition Lakes 7.0 pH 6.0- 5.0- 4.0; Acid Lakes \ .............. : .:,:,: .:.:,:,::: _ ::::::::::::::::::::::::::: _ ............... ~ .............. :t: :-::: :-:-:-:-: :-: .,\....................... .-.~-.- - -. ::~:::::::::: :: i. . - -: :-:2:2:, ·:-:-:l2.2.- 2,- ·.,..2,~..-,..:..,2-.. :-::-:-::~::::: . ~ \ -2-.2.- -. :-:-: . . . 100 H ~ added, ~eq/l -100 H CO3—peq/1 - 50 200 FIGURE 8.5 Titration curve for bicarbonate solution at a concen- tration of 100 ,ueq/lj illustrating the acidification process. SOURCE: Henriksen (1980).

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149 acid forms, there is a potential for ecosystem acidification by the Vitrification of ammonia, from atmospheric precipitation or from the decomposition of dead organic matter. Two equivalents of hydrogen ion are generated for each equivalent of ammonium ion transformed to NO3 (Reuse 1975a), but one of those equivalents may be consumed upon either uptake or denitrification by components of the biota. If all of the ammonium and nitrate ions are transformed and/or utilized by organisms, the net potential for ecosystem acidification will be measured by the difference of ammonia and nitrate in equivalents. Preliminary data show very high concentrations of ammonia in precipitation over the central United States, possibly as a result of crop fertilization and livestock culture (Figure 8.6~. Oden (1976) has compared the trends in southern Sweden of anthropogenic acidification directly by mineral acids and indirectly by biological processes, and the normal background of biological acidification by nitrogen transformation (Figure 8.71. Mayer (1979) discussed the conditions under which natural acidification may be most significant. Extent of the Problem Large areas of the earth's surface consist of poorly buffered geologic materials. Where such areas occur within several hundred kilometers of sources of atmospheric emissions that are acid precursors, detectable acidification of at least freshwater ecosystems may be expected to occur. In the united States such areas occur largely in the eastern part of the country (Cogbill and Likens 1974~. In Canada, there are roughly 2 million square kilometers of acid-sensitive terrain--most of the eastern half of the country. Similar large geologic areas occur in Scandanivia, Scotland, and the northern part of the Soviet Union. A high proportion of the world's freshwaters occur in such terrain (between 50 and 80 percent depending on estimates); thus, large-scale degradation becomes an important concern. EFFECTS OF AC ID ON THE B IOS PHERE Aquatic Ecosystems Widespread and pronounced effects of acid precipitation have been recorded in poorly buffered aquatic ecosystems. The low chemical capacity of such "soft" waters to buffer against increasing hydrogen-ion supply causes the water to be in precarious pH balance even under natural conditions. Indeed, many natural lakes may become more acidic over hundreds or thousands of years owing to drainage from acid peat bogs. The bog moss, Sphagnum, generates in its cell walls polyuronic acids (Clymo 1964), whose hydrogen ions are exchanged for metal cations from precipitation (Gorham and Cragg 1960~. The hydrogen ions that are released in this way acidify the bog waters.

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172 The soils most susceptible to rapid acidification are well-drained brown forest soils (alfisols) that are sandy and noncalcareous but not already strongly acid (Wiklander 1973, 1974, 1979~. Such coarse soils are moderately to highly saturated by "basic" cations such as Ca++, Mg++, K+, and Na+, which are leached away as they are replaced by "acid" cations such as hydrogen ions, aluminum ions, and hydroxy-aluminum ions. As the exchange complex becomes dominated increasingly by the "acid" cations, the pH of the soil declines. Fine-textured mineral soils, because of their high clay content, have a much greater cation-exchange capacity, which is usually strongly saturated by "basic" cations, and such soils are much better buffered against acidification. Strongly acid podzol (spodosol) soil horizons--whether organic, with a high cation-exchange capacity, or sandy, with low cation-exchange capacity--are only slightly susceptible to further acidification. Their exchange sites are already dominated by the "acid" cations, and thus added hydrogen ions are more likely to pass through the system in the percolating waters. Nevertheless, because such soils are already impoverished, even slight losses owing to increased acidification may be critical for soil fertility. The influence of acid rain upon the impoverished lateritic soils of the tropics could also be very significant but does not seem to have been examined. An interesting problem is the synergistic influence of neutral salts upon cation exchange under acid conditions. Wiklander (1975, 1979, cf. Abrahamsen et al. 1979) has made the important point that adding divalent neutral sulfates--often abundant in acid rain--to acid leaching solutions may retard significantly the replacement and loss of basic cations from already acid soils and thus favor the acidification of receiving waters. Monovalent chlorides have a lesser effect. Addition of strong acids to precipitation will cause not only ion exchange on the surfaces of soil particles but also alteration of the particles themselves by weathering, either directly or by increasing the hydrogen-ion saturation of organic and inorganic soil colloids, which act as weathering agents (Loughnan 1969~. Carbonates are weathered very readily indeed, but unless they are present only in small amounts (cf. Salisbury 1922, 1925) effects upon the soil are likely to be extremely slight. Aluminosilicate minerals are more slowly dissolved by acid rain, but according to Norton (1976, cf. Johnson 1979) the solubility of aluminum compounds such as gibbsite, amorphous aluminum hydroxide, and kaolinite increases rapidly below a pH of about 505. In this connection, lake waters generally show an order-of-magnitude rise in dissolved aluminum as pH falls from 5.5 to around 4 (Wright et al. 1976~. The solubility of hydrous oxides of iron, produced by the weathering of iron-bearing minerals, is affected by a lowering of pH to a somewhat lesser degree than oxides of aluminum, according to Black (1967, see also Birkeland 1974~. The dissolution of silica is essentially unaffected over the pH range 2 to 8 but rises rapidly above pH 8 (Birkeland 1974~.

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173 Increased leaching of potassium, calcium, aluminum, and magnesium due to ion exchange reactions with hydrogen ions is one of the most commonly observed effects of acid precipitation (Figure 8.19; also see review by Abrahamsen 1980~. These increased rates of leaching appear to outstrip compensatory increases in weathering, reducing the exchangeable pool of the above cations. Sulfate is generally adsorbed in soil, largely in the B horizon (Farrell et al. 1980, Singh 1980, Figure 8.20~. Whitby and Hutchinson (1974) found that soils acidified to pH values below 4.0 by smelter fumigations in the Sudbury, Ontario area released sufficient aluminum into the soil solution to severely inhibit the establishment and elongation of seedling roots. Ulrich et al. (1980) have suggested that elevated aluminum concentrations caused the crown dieback in beech (Fagus sylvatica) and the failure of seedlings in the Solling project research forest in Germany. The direct role of aluminum in this response as opposed to the effect of severe droughts which occurred during the same period has not been established. It is not clear at present whether the effects of acid precipitation on soil will result in long-term degradation of terrestrial ecosystems. At least one controlled 5-year study has revealed a significant depletion of cation exchange capacity in soils subjected to loading of 13.3 k. eq. H+ per ha of artificial acid rain (Farrell et al. 1980~. Troedsson (1980) found declining quantities of Ca++, Mg++, and Kit in Swedish soils over a 10-year period. Both theoretical and experimental studies suggest that such declines may be widespread in vulnerable soils after several decades to a few centuries of acid precipitation (Oden 1968, Reuss 1975b, Maimer 1976, Norton 1976, Tamm 1977, McFee 1978, Abrahamsen and Stuanes 1980~. Differences in effects of different levels of acidification upon trace metals in soils are even less well understood. The only detailed study is that of Tyler (1978), in which purely organic mor humus layers (Romell 1932, 1935, Lutz and Chandler 1946) from Swedish spruce forests were leached in the laboratory by simulated rain acidified to 5 different pH levels. Two sets of humus layers were examined, one far from and the other near to a brass mill. The latter exhibited very strong contamination by copper (695 x the control set, far from the pollution source) and zinc (124 x), with lesser enrichment of cadmium (24 x), lead (7.6 x), chromium (3.7 x), manganese (3.1 x), nickel (2.2 x), and vanadium (1.3 x). Simulated rains at pH values of 4.2, 3.4, 3.2, 3.0, and 2.8 were applied over 125 days in amounts totaling 625 ml/g of humus. The s~mulated rain at pH 4.2, a value often reached in Swedish precipitation, released (over 125 days) substantial percentages of several trace metals in the control humus layers, notably manganese (44~), nickel (44%), cadmium (33%), and zinc (25~. Percentage release was much less for the other metals, copper (12~), vanadium (11~), chromium (8.9~), and lead (2.2~. With the notable exception of vanadium release (40~), percentage releases were generally less in the humus layers of the polluted site, as follows: manganese (4.1~), nickel (11%), cadmium (9.3~), zinc (18%), copper (1.7%), chromium

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174 pH 2 1400 - 800- 700 600 N - ~ 500 a) 400 300 - 200- 100 - O- -100 ~ pH4 oH 3 control, not watered Ca Mg K I: K,Mg,Ca ~ _ ~ . us SOURCE: Stuane s ~ l 98 0) FIGURE 8.19 Loss of nutrients by weathering and mineralization. Artificial rain of varying acidity was applied in a field experiment over a 5-year period. The "rain" amounted to 50 mm per month for the 5 frost-free months of the year. SOURCE: Stuanes (1980).

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175 35So2- S adsorbed 3 0 25 50 75 100 125 1 1 1 1 ' 1 E ~ P Bsl Bs2 Bc . C1 C2 C3 23 30 50 72 Q 95 a, . _ o cat . \\\\\\\\\\\\\\\\\\\\\\\\~ ~~\~\~: \\\\ I ~r- 8 E 25 Bs 76 2C 134 142 - 0 25 50 75 100 0 25 n ~ ~~ 6 Or 7 Ah Bw 20W 4 Ah 63 2C 90 0 25 50 75 ~ I ~ ~~ P. \\\\\\\\\\\ 4 - . ~ 50 0 15°07 A2BhC2r FIGURE 8.20 Distribution of adsorbed 3sSO4 sulfur in soil profiles. PI and Ps are iron-podzols; P4 iS a semipodzol (all typic udipsamments); and P6 and P7 are brown- earths (umbric dystrochrept and aquic haploboroll, respectively). SOURCE: Singh (1980). 25 50 ' 1 1 P7

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176 (1.5~), and lead (0.24%~. This difference is most likely due to the much lesser acidification of the polluted humus, which yielded a percolate with a pH of 6.1 at the end of the experiment, in contrast to a pH of 4.5 in the percolate from the unpolluted humus. The effects of acidification at 5 different pH levels in this experiment may be examined by comparing releases at each of the more acid pH values to that at pH 4.2, taking the release at this pH as unity. Figure 8.21 demonstrates that once again no simple rule can be given. For instance, increasing acidification has much less effect upon lead than upon zinc at all pH values above 2.8, but at pH 2.8 the release of lead is strongly accentuated and exceeds that of zinc in both polluted and unpolluted soils. In the case of vanadium, acidification reduces the amount released, except from the unpolluted humus at the lowest pH, 2.8, from which the release is doubled. In the polluted humus, increasing acidification below pH 4.2 has about the same effect upon zinc and nickel, whereas in the unpolluted humus, acidification has a much greater effect upon zinc. Elucidation of the influence of acid precipitation upon the release of polyvalent metals from mineral soils will be greatly complicated by the fact that organic acids and polyphenols produced by organisms are also of much importance in the mobilization of oxides of aluminum, iron, and manganese from soils, and breakdown of these oxides will release the many trace metals adsorbed by them (several references in Russell 1973~. Presumably acid rain--with its strong mineral acids--will have some effect upon the weathering action of the organic acids and polyphenols, the latter being known to reduce iron more strongly in acid than in neutral conditions (Russell 1973~. Acidification can also reduce the stability of the fulvic acid components of humic acids in soils and their metal complexes, and thus metal availability should be greater at lower pH (Schnitzer 1980~. Little is known about such interactions, however, and they should be given greater attention. A number of soil microbial processes appear to be affected by acidification. Reduced soil pH may cause a reduction in nitrogen fixation (Alexander 1980), although mineralization of organic nitrogen may increase (Nyborg and Hoyt 1978~. Increased acidity also causes Vitrification to decrease; vitrification usually ceases entirely at about pH 4.0. There may be some degree of acclimation in acid environments (Walker and Wickramasinghe 1979~. Denitrification, the main means by which nitrogen is released to the atmosphere from the biosphere, is also affected by acidification (Figure 8.22~. Decreased rates, and a change in end product from molecular nitrogen to nitrous oxide, are observed in anoxic environments at pH values less than 6.0. At lower pH levels an appreciable yield of nitric oxide is observed, which is phytotoxic (Wijler and Delwiche 19541. Only the mechanisms involved in nitrogen fixation have been studied directly. Legumes, which depend on Rhizobium for nitrogen fixation, are particularly sensitive to acidification. Alexander (1980) suggests that this may be due to the high concentrations of A1, Mn, or Fe in acid soils. Another possibility is limitation by

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177 U NPOLLUTED SOIL POLLUTED SOIL 4 3 2 o LLJ 1 - \ Pb _ ~ \ ~ _ O ~: o 4 z o 3 cr 2 3 _ 2 _ 1 _ O _ C U ~_ ' ' ' ' , ~ 2.8 3.0 3.2 3.4 4.2 , I I l , ~ 2.8 3.0 3. 2 3.4 4.2 I ' ' ' ,' ' 2.8 3.0 3.2 3.4 4.2 4C 30 20 10 o Cul I Pb 1~1 I T I I ~ 2.8 3.0 3.2 3.4 4.2 pH OF LEACHING SOLUTION FIGURE 8.21 Effect of acidity upon the leaching of heavy metals from spruce humus layers, normalized to pH 4.2. Left-hand graphs represent unpolluted humus layers; right- hand graph represents humus layers polluted from a brass mill.

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I78 00 50 40 30 20 10 o V= 12.Sx-47.2 r= 0.~ , i./ 3 4 @ e 6 7 pH FIGURE 8.22 CoIIelation Bitten denitdOcation Ivies and sod pH. SOURCE: ~UDeI ~ at. (1980t

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179 molybdenum, an essential element for nitrogen fixation in legumes, because Mo solubility decreases at low pH. Francis et al. (1980) found that increased acidity decreased rates of decomposition, ammonification, Vitrification, denitrification, and N2 fixation in soil. Microbial degradation of pesticides was also reduced at low pH. Lohm (1980) found that the biomass of fungal mycelium increased in artifically acidified forest soils, replacing bacteria, which were reduced in number under acid conditions. Springtails also increased. The number of Enchytraeidae (oligochaete worms), Collembola (spring/ails), and Acari (mites) also decreased under acid conditions. The overall result was a net decrease in decomposition. Haagvar's (1980) study of invertebrates under acid conditions produced analogous results. It appears that the detrimental effects of acidification on soils may not be due to--or at least, wholly attributable to--acidification per _ but may be caused by the altered concentrations of nutrients and heavy metals in acidified soils. Direct toxicity of hydrogen ions in soils was found to be negligible over the pH range 4 to 8 (Ar non and Johnson 1942, Arnon et al. 1942~; around pH 3 roots may be injured and above pH 8 phosphate absorption may be inhibited. In general, pH affects plant growth--and thereby the cycles of many (especially biophile) elements--by influencing the concentrations of different ions in the soil solution (Russell 1973, Nyborg 1978, Hutchinson and Collins 1978~. Some of the elements affected are major plant nutrients (e.g., calcium, potassium, phosphate, nitrate), and others are toxicants (e.g., aluminum, lead), while several trace elements (e.g., manganese, copper, zinc) may be nutrients at low concentrations and toxicants at high concentrations (Bowen 1966), as for example around metal smelters (Stokes et al. 1973, Hutchinson and Whitby 19771. The biogeochemistry of an ecosystem varies systematically as both vegetation and soil change in the course of succession, and acidification (whether normal or anthropogenic) plays a marked role in this process (Gorham et al. 1979~. For a time, the leaching effect of acid precipitation upon a soil with appreciable reserves of calcium may enrich the stream and lake waters draining such soils without acidifying them (Gordon and Gorham 1963, Gorham 1978a), because the acids will be neutralized in percolating through the upland soils and liberating basic cations. However, if water flow occurs mainly as surface runoff, for example over frozen ground in spring, or downward through old root channels, animal burrows, and along rock faces (Tamm and Troeds son 1957~--then acidification of the receiving water may take place even if the upland soil possesses quite substantial buffering capacity. As time goes on and upland soils become more acid, they also become less susceptible to further acidification by acid rain, because exchange sites within the soil are already highly saturated by hydrogen ions. In this situation, substantial amounts of acid will percolate to the receiving waters of streams and lakes, which consequently undergo a pronounced decline in pH. Acidification may be regarded as a normal tendency of ecosystem succession on base-poor substrata, because the biota produce acids of

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180 various kinds metabolically (carbonic acid and various organic acids through the oxidation of organic carbon compounds; nitric acid by Vitrifying bacteria; sulfuric acid by bacterial oxidation of organic sulfur compounds; and hydrogen ions attached to the surfaces of roots, Sphagnum mosses, fungal hyphae, and bacteria, cf. Wiklander 1979~. But natural acidification is accelerated appreciably by acid rain, which often brings about a change within decades in vulnerable aquatic ecosystems, once their buffering capacity is exhausted. Acid rain may also cause a slower change--perhaps over centuries--in upland ecosystems, where soil minerals and ion-exchange complexes provide a greater degree of buffering, even in the most vulnerable watersheds. If acidification should lead to substantial reduction in soil base saturation, recovery following removal of the acid loading could well take decades to centuries. Both natural and anthropogenic acidification processes in terrestrial soils deserve increased study, so that we may assess their relative importance in different ecosystems. Wetlands The disappearance of several species of the bog moss Sphagnum from the vast blanket bogs of the southern Pennines in the British Isles is the major vegetational change that has resulted from atmospheric pollution in that country since the Industrial Revolution (Talks 1964~. In this connection, Gorham (1958b) has shown that pools in bogs with an intact Sphagnum cover ranged in pH from 4.5 in remote areas of Britain, where the acidity is chiefly biogenic, to as low as 3.9 in the northern Pennines closer to urban/industrial centers. In these bog waters the correlation between HE and non-marine SO4 ions was highly significant (r = 0.985~. Near Sheffield in the southern Pennines the pH of bog pools was only 3.25, and the concentration of SO4 reached 46 mg/1. Although it is impossible to ascribe with certainty the disappearance of the bog moss from the southern Pennines to acid precipitation, recent experiments with the same Sphagnum species indicate that either acid rain or SO2 fumigation alone have detrimental effects consistent with the observed disappearance (Ferguson, Lee, and Bell 1978~. In contrast, Hemond (1980) has calculated that in a Massachusetts peatland dominated by Sphagnum the reduction of sulfate and biological utilization of nitrate totally buffer the effect of acid deposition upon interstitial waters there. Vast areas of sphagnum bog occur in North America, including portions of northern Minnesota and the lowlands of Hudson and James bays where rain is now quite acidic. Many of these peatland ecosystems have very low buffering capacities, which suggests that they may be easily unbalanced by acid precipitation. Although data

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181 are sparse, the abundance of peatlands in northern latitudes around the globe suggests that they may be significant reservoirs in the global cycles of many elements. ASSESSMENT OF ECOLOGICAL EFFECTS For improvements in our understanding of the effects of acid deposition we must concentrate on two major research areas. First, we need long-term monitoring (Botkin 1978) of sensitive organisms or communities (e.g., the "neuston" of freshwater surfaces, cf. Gorham 1976, 1978a,b) in especially vulnerable ecosystems, so that we can receive early warning of ecosystems at hazard. Such monitoring presupposes the identification of a series of indicator organisms and communities (Thomas 1972) or indices to community structure (Cairns 1974) and the development of an adequate scheme for rating ecosystem vulnerability. The Calcite Saturation Index developed by Conroy et al. (1974, see also Kramer 1976) provides a useful guide to the vulnerability of lakes, but because it uses only the concentrations of calcium, bicarbonate, and hydrogen ions in the water, it does not take into account the full range of factors involved (cf. Table 8.11. Second, we need experimental ecosystem-scale studies to examine linkages among and processes within uplands, wetlands, streams, and lakes (Likens and Bormann 1974, Schindler 1980b) and the mechanisms by which acid deposition alters ecosystem function. Some of these studies could be conducted on watersheds exposed to ambient levels of acid rain, while other studies, in relatively unpolluted areas, could subject whole ecosystems to experimental acidification. Perhaps the recently proposed national network of experimental ecological reserves (TIE 1977) could provide suitable sites for such studies, as could several of the experimental watersheds (Table 22 in Likens et al. 1977) set up for other purposes. One of these, the Hubbard Brook Experimental Forest, has already become an important site for research on acid rain (Likens et al. 1977~. AMELIORATION Of the options presently available only the control of emissions of sulfur and nitrogen oxides can significantly reduce the rate of deterioration of sensitive freshwater ecosystems. It is desirable to have precipitation with pH values no lower than 4.6 to 4.7 throughout such areas, the value at which rates of degradation are detectable by current survey methods, as mentioned above. In the most seriously affected areas (average precipitation pH of 4.1 to 4.2), this would mean a reduction of 50 percent in deposited hydrogen ions. Control of SO2 from new electrical generating plants alone would be insufficient to accomplish this, and thus restrictions on older plants must be considered. Furthermore, there are no proposed restrictions on the emission of nitrogen oxides, and the amounts of these substances emitted are expected to continue to increase (see Figure 4.6).

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182 The alkalinity of waters endangered by acidification can be enhanced by a number of means, most notably by "liming"--adding calcium carbonate or oxide--and by adding phosphorus to stimulate biological fixation of nitrate and CO2. All of these techniques are expensive ($50 and more per hectare of water surface), and treatments must be repeated every few years. Due to high costs and logistic difficulties, lime cannot be applied to the vast areas that are currently endangered by acidification. In the areas most susceptible to acid deposition, it will therefore be impossible to maintain the alkalinity and pH of more than a few selected bodies of water. Furthermore, pH and alkalinity cannot be artificially maintained without increasing the ionic concentration of the receiving water, the consequences of which have not been investigated. The fate of dissolved toxic metals after liming is also poorly known. Addition of nutrients to increase alkalinity has been investigated at a few sites in Ontario, both by the Ontario Ministry of Environment and Canadian Department of Fisheries and Oceans, but results are not available yet. Many of the objections to liming will also apply to nutrient additions. SUMMARY Acid deposition, due to the further oxidation of sulfur and nitrogen oxides released to the atmosphere by anthropogenic sources, is causing widespread damage to aquatic ecosystems, including loss of bicarbonate, increased acidity, and higher concentrations of toxic metals. As a result, several important species of fish and invertebrates have been eliminated over substantial parts of their natural ranges. Effects on terrestrial ecosystems are less pronounced. Increased leaching of both nutrients and toxic elements is evident in poorly buffered soils sensitive to acidification. There is some evidence for damage to crop plants, and many soil microbial processes are negatively affected at low pH. Trees appear to be slightly stimulated by acid precipitation, although this effect is expected to be short-lived, because of increased leaching of cationic nutrients and the buildup of toxic co wentrations of metals in soil water. Better long-term studies of deposition processes and of effects on ecosystems are required to illuminate the complex ecological effects of acid precipitation and associated nutrients and toxicants. The control of emissions of sulfur and nitrogen oxides from fossil fuels is necessary to halt the acidification of sensitive aquatic ecosystems.