Acid rain was discovered in the 19th century by Robert Angus Smith, a pharmacist from Manchester (England), who measured high levels of acidity in rain falling over industrial regions of England and contrasted them to the much lower levels he observed in less polluted areas near the coast. Little attention was paid to his work until the 1950s, when biologists noticed an alarming decline of fish populations in the lakes of southern Norway and traced the problem to acid rain. Similar findings were made in the 1960s in North America (the Adirondacks, Ontario, Quebec). These findings spurred intense research to understand the origin of the acid rain phenomenon.
Rainwater falling in the atmosphere always contains impurities, even in the absence of human influence. It equilibrates with atmospheric CO2, a weak acid, following the reactions presented in chapter 6:
The corresponding equilibrium constants in dilute solution at 298 K are KH = [CO2H2O]/PCO2 = 3x10-2 M atm-1, K1 = [HCO3-][H+]/[CO2H2O] = 4.3x10-7 M (pK1 = 6.4), and K2 = [CO32-][H+]/[HCO3-] = 4.7x10-11M (pK2 = 10.3). From these constants and a preindustrial CO2 concentration of 280 ppmv one calculates a rainwater pH of 5.7. Other natural acids present in the atmosphere include organic acids emitted by the biosphere, HNO3 produced by atmospheric oxidation of NOx originating from lightning, soils, and fires ( section 11.4 ), and H2SO4 produced by atmospheric oxidation of reduced sulfur gases emitted by volcanoes and by the biosphere. A comparative analysis of these different natural sources of acidity is conducted in See Acid rain: the preindustrial atmosphere . The natural acidity of rain is partly balanced by natural bases present in the atmosphere, including NH3 emitted by the biosphere and CaCO3 from suspended soil dust.
When all of these influences are taken into account, the pH of natural rain is found to be in the range from 5 to 7. The term acid rain is customarily applied to precipitation with a pH below 5. Such low pH values are generally possible only in the presence of large amounts of anthropogenic pollution.
Figure 13-1 shows the mean pH values of precipitation measured over North America. pH values less than 5 are observed over the eastern half. We can determine the form of this acidity by examining the ionic composition of the precipitation; data for two typical sites are shown in See Median concentrations of ions (meq l-1) in precipitation at two typical sites in the United States. . For any precipitation sample, the sum of concentrations of anions measured in units of charge equivalents per liter must equal the sum of concentrations of cations, since the ions originated from the dissociation of neutral molecules. This charge balance is roughly satisfied for the data in See Median concentrations of ions (meq l-1) in precipitation at two typical sites in the United States. ; an exact balance would not be expected because the concentrations in the Table are given as medians over many samples.
Consider the data for the New York site in See Median concentrations of ions (meq l-1) in precipitation at two typical sites in the United States. . The median pH at that site is 4.34, typical of acid rain in the northeastern United States. The H+ ion is the dominant cation and is largely balanced by SO42- and NO3-, which are the dominant anions. We conclude that H2SO4 and HNO3 are the dominant contributors to the precipitation acidity. Both are strong acids which dissociate quantitatively in water to release H+:
As shown in Figure 13-1 , SO42- and NO3- concentrations throughout the United States are more than enough to balance the local H+ concentrations. More generally, analyses of rain composition in all industrial regions of the world demonstrate that H2SO4 and HNO3 are the main components of acid rain.
Consider now the data for the southwest Minnesota site in See Median concentrations of ions (meq l-1) in precipitation at two typical sites in the United States. . The concentrations of SO42- and NO3- are comparable to those of the New York site, indicating similar inputs of H2SO4 and HNO3. However, the H+ concentration is two orders of magnitude less; the pH is close to neutral. There must be bases neutralizing the acidity. To identify these bases, we examine which cations in See Median concentrations of ions (meq l-1) in precipitation at two typical sites in the United States. replace the H+ originally supplied by dissociation of HNO3 and H2SO4. The principal cations are NH4+, Ca2+, and Na+, indicating the presence in the atmosphere of ammonia (NH3) and alkaline soil dust ( CaCO3, Na2CO3). Ammonia dissolved in rainwater scavenges H+:
The equilibrium constant for (R7) is K = [NH4+]/[NH3(aq)][H+] = 1.6x109 M, so that[NH4+]/[NH3(aq)] = 1 for pH = 9.2. At the pH values found in rain, NH3 behaves as a strong base; it scavenges H+ ions quantitatively and NH4+ appears as the cation replacing H+. Neutralization of H+ by dissolved soil dust proceeds similarly:
The relatively high pH of rain in the central United States ( Figure 13-1 ) reflects the large amounts of NH3 emitted from agricultural activities (fertilizer use, livestock), and the facile suspension of soil dust due to the semi-arid climate. Note from Figure 13-1 that NH4+concentrations are maximum over the central United States.
It has been known since the 1960s that the high concentrations of HNO3 and H2SO4 in acid rain are due to atmospheric oxidation of NOx and SO2 emitted by fossil fuel combustion. Understanding of the oxidation mechanisms is more recent. The mechanisms for oxidation of NOx to HNO3 were discussed in chapters 10 and 11, and in See Nighttime oxidation of NOx . Both OH (day) and O3 (night) are important oxidants and lead to a NOx lifetime over the United States of less than a day. We focus here on the mechanisms for oxidation of SO2 to H2SO4.
Sulfur dioxide (SO2) is emitted from the combustion of sulfur-containing fuels (coal and oil) and from the smelting of sulfur-containing ores (mostly copper, lead, and zinc). In the atmosphere, SO2 is oxidized by OH to produce H2SO4:
The lifetime of SO2 against reaction with OH is 1-2 weeks. A major research problem in the 1970s was to reconcile this relatively long lifetime with the observation that SO42- concentrations in rain are maximum over SO2 source regions ( Figure 13-1 ). This observation implies that SO2 must be oxidized to H2SO4 rapidly; otherwise, the emission plume would be transported far from the SO2 source region by the time it was oxidized to H2SO4. Research in the early 1980s showed that most of the atmospheric oxidation of SO2 actually takes place in cloud droplets and in the raindrops themselves, where SO2 dissociates to HSO3- which is then rapidly oxidized in the liquid phase by H2O2 produced from self-reaction of HO2 (chapter 11):
Reaction (R16) is acid-catalyzed (note the presence of H+ on the left-hand side). The rate of aqueous-phase sulfate formation is
where the K's are equilibrium constants. Acid catalysis is key to the importance of (R16) for generating acid rain; otherwise, the reaction would be suppressed at low pH because [HSO3-] depends inversely on [H+] by equilibrium (R14) (see See Aqueous-phase oxidation of SO2 by ozone ). Substituting numerical values in (13.1) indicates that reaction (R16) is extremely fast and results in titration of either SO2 or H2O2 in a cloud; measurements in clouds show indeed that SO2 and H2O2 do not coexist. Reaction (R16) is now thought to provide the dominant atmospheric pathway for oxidation of SO2 to H2SO4, although there are still unresolved issues regarding the mechanism for oxidation during the winter months when production of H2O2 is low.
Acid rain falling over most of the world has little environmental effect on the biosphere because it is rapidly neutralized after it falls. In particular, acid rain falling over the oceans is rapidly neutralized by the large supply of CO32- ions (chapter 6). Acid rain falling over regions with alkaline soils or rocks is quickly neutralized by reactions such as (R9) taking place once the acid has deposited to the surface. Only in continental areas with little acid-neutralizing capacity is the biosphere sensitive to acid rain. Over North America these areas include New England, eastern Canada, and mountainous regions, which have granitic bedrock and thin soils ( Figure 13-2 ).
In areas where the biosphere is sensitive to acid rain, there has been ample evidence of the negative effects of acid rain on freshwater ecosystems. Elevated acidity in a lake or river is directly harmful to fish because it corrodes the organic gill material and attacks the calcium carbonate skeleton. In addition, the acidity dissolves toxic metals such as aluminum from the sediments. There is also ample evidence that acid rain is harmful to terrestrial vegetation, mostly because it leaches nutrients such as potassium and allows them to exit the ecosystem by runoff. Although NH3 in the atmosphere neutralizes rain acidity by formation of NH4+, this acidity may be recovered in soil when NH4+ is assimilated into the biosphere as NH3 or goes through the microbial nitrification process ( section 6.3 ):
Beyond the input of acidity, deposition of NH4+ and NO3- fertilizes ecosystems by providing a source of directly assimilable nitrogen ( section 6.3 ). This source has been blamed as an important contributor to the eutrophication (excess fertilization) of the Chesapeake Bay; a consequence of this eutrophication is the accumulation of algae at the surface of the bay, suppressing the supply of O2 to the deep-water biosphere. Recent studies of terrestrial ecosystems in the United States show that increases in NH4+ and NO3- deposition do not stimulate growth of vegetation but lead instead to accumulation of organic nitrogen in soil. The long-term implications of this nitrogen storage in soil are unclear.
Emissions of SO2 and NOx in the United States and other industrial countries have increased considerably over the past century due to fossil fuel combustion. Since 1970, however, emissions in the United States and Europe have leveled off due to pollution control efforts. Emissions of SO2 in the United States have decreased by 25% since 1970, while NOx emissions have remained flat. Technology for SO2 emission control is expensive but readily available (scrubbers on combustion stacks, sulfur recovery during oil refining). The control of SO2 emissions in the United States was initially motivated by the air quality standard for SO2 (SO2 is a toxic gas) rather than by concern over acid rain; the original Clean Air Act of 1970 did not include acid rain under its purview. The revised Clean Air Act, which now targets acid rain, mandates a further decrease by a factor of 2 in SO2 emissions from the United States over the next decade. Similar steps to further decrease SO2 emissions in the future are being taken in European countries. By contrast, SO2 emissions in eastern Asia are on a rapid rise fueled in large part by the industrialization of China and India relying on coal combustion as a source of energy. A serious acid rain problem may develop over eastern Asia in the decades ahead.
In the United States, the SO2 control measures to be achieved under the revised Clean Air Act will provide significant environmental relief over the next decades but cannot be expected to solve the acid rain problem. First, current forecasts indicate little decrease in NOx emissions over the next decade. Second, a reduction of acid levels by a factor of 2 is not enough to make rain innocuous to the biosphere (note that decreasing [H+] by a factor of 2 increases the pH by only 0.3 units). Third, acid rain is in part a cumulative problem; as the acid-neutralizing capacity of soils gets depleted the ecosystems become increasingly sensitive to additional acid inputs.
Recently, there has been considerable interest in the possibility that reductions of SO2 emissions to combat the acid rain problem might have negative side effects on climate. As discussed in section 8.2.3 , there is evidence that anthropogenic sulfate aerosol at northern midlatitudes has caused regional cooling, perhaps compensating in a complicated way for the effect of greenhouse warming. Could reductions in SO2 emissions expose us to the full force of global warming? This is an interesting research question which should not, however, discourage us from going ahead with SO2 emission reductions. We do not understand climate well enough to play with radiative forcing effects of opposite sign and hope that they cancel each other.