The Nature And Development Of Acid Rain

Acid rain is normally considered to be a byproduct of modern atmospheric pollution. Even in a pure, uncontaminated world, however, it is likely that the rainfall would be acidic. The absorption of carbon dioxide by atmospheric water produces weak carbonic acid, and nitric acid may be created during thunderstorms, which provide sufficient energy for the synthesis of oxides of nitrogen (NOX) from atmospheric oxygen and nitrogen. During volcanic eruptions or forest fires, sulphur dioxide (SO2) is released into the atmosphere to provide the essential component for the creation of sulphuric acid. Phytoplankton in the oceans also emit sulphur during their seasonal bloom period. The sulphur takes the form of dimethyl sulphide (DMS)

which is oxidized into SO2 and methane sulphonic acid (MSA). The MSA is ultimately converted into sulphate (Cocks and Kallend 1988). Acids formed in this way fall out of the atmosphere in rain to become involved in a variety of physical and biological processes once they reach the earth's surface. The return of nitrogen and sulphur to the soil in naturally acid rain helps to maintain nutrient levels, for example. The peculiar landscapes of limestone areas—characterized by highly weathered bedrock, rivers flowing in steep-sided gorges or through inter-connected systems of underground stream channels and caves—provide excellent examples of what even moderately acid rain can do.

In reality, since 'acid rain' includes snow, hail and fog as well as rain, it would be more appropriate to describe it as 'acid precipitation'. The term 'acid rain' is most commonly used for all types of 'wet deposition', however. A related process is 'dry deposition', which involves the fallout of the oxides of sulphur and nitrogen from the atmosphere, either as dry gases or adsorbed on other aerosols such as soot or fly ash (Park 1987). As much as two-thirds of the acid precipitation over Britain falls as dry deposition in the form of gases and small particles (Mason 1990). On contact with moisture in the form of fog, dew or surface water they produce the same effects as the constituents of wet deposition. At present, both wet and dry deposition are normally included in the term 'acid rain' and, to maintain continuity, that convention will be followed here.

Current concern over acid rain is not with the naturally produced variety, but rather with that which results from modern industrial activity. Technological advancement in a society often depends upon the availability of metallic ores, which can be smelted to produce the great volume and variety of metals needed for industrial and socio-economic development. Considerable amounts of SO2 are released into the atmosphere as a by-product of the smelting process, particularly when non-ferrous ores are involved. The burning of coal and oil, to provide energy for space heating or to fuel thermal electric power stations, also produces SO2. The continuing growth of transportation systems using the internal combustion engine—another characteristic of a modern technological society— contributes to acid rain through the release of NOX into the atmosphere.

Initially, the effects of these pollutants were restricted to the local areas in which they originated, and where their impact was often obvious. The detrimental effects of SO2 on vegetation around the smelters at Sudbury (Ontario), Trail (British Columbia), Anaconda (Montana) and Sheffield (England) have long been recognized, for example (Garnett 1967; Hepting 1971). As emissions increased, and the gases were gradually incorporated into the larger scale atmospheric circulation, the stage was set for an intensification of the problem. Sulphur compounds of anthropogenic origin are now blamed for as much as 65 per cent of the acid rain in eastern North America, with nitrogen compounds accounting for the remainder (Ontario: Ministry of the Environment 1980). In Europe, emission totals for SO2 and NOX are commonly considered to split closer to 75 per cent and 25 per cent (Park 1987). Since the early 1970s, however, declining SO2 emissions and a growing output of NOX have combined to bring

Scale Showing Acid Rain

Figure 4.1 The pH scale: showing the pH level of acid rain in comparison to that of other common substances

Figure 4.1 The pH scale: showing the pH level of acid rain in comparison to that of other common substances

Figure 4.2 Schematic representation of the formation, distribution and impact of acid rain

Figure 4.2 Schematic representation of the formation, distribution and impact of acid rain

Formation Acid Rain
Source: Compiled from information in Park (1987); Miller (1984); LaBastille (1981)

the relative proportions of these gases close to the North American values (Mason 1990).

Acid precipitation produced by human activities differs from natural acid precipitation not only in its origins, but also in its quality. Anthropogenically produced acid rain tends to be many times more acidic than the natural variety, for example. The acidity of a solution is indicated by its hydrogen ion concentration or pH (potential hydrogen); the lower the pH, the higher the acidity. A chemically neutral solution, such as distilled water, has a value of 7, with increasingly alkaline solutions ranging from 7 to 14, and increasingly acidic solutions ranging from 7 down to 0 (see Figure 4.1). Since this pH scale is logarithmic, a change of one point represents a tenfold increase or decrease in the hydrogen ion concentration, while a two-point change represents a one hundredfold increase or decrease. A solution with a pH of 4.0 is ten times more acidic than one of pH 5.0; a solution of pH 3.0 is one hundred times more acidic than one of pH 5.0.

The difference between 'normal' and 'acid' rain is commonly of the order of 1.0 to 1.5 points. In North America, for example, naturally acid rain has a pH of about 5.6, whereas measurements of rain falling in southern Ontario, Canada, frequently provide values in the range of 4.5 to 4.0 (Ontario: Ministry of the Environment 1980). To put these values in perspective it should be noted that vinegar has a pH of 2.7 and milk a pH of 6.6 (see Figure 4.1). Thus, Ontario rain is about 100 times more acidic than milk, but 100 times less acidic than vinegar. Similar values for background levels of acidic rain are indicated by studies in Europe. The Central Electricity Generating Board (CEGB) in Britain has argued for pH 5.0 as the normal level for naturally acid rain (Park 1987), but the average annual pH of rain over Britain between 1978 and 1980 was between 4.5 and 4.2 (Mason

1990). Remarkably high levels of acidity have been recorded on a number of occasions on both sides of the Atlantic. In April 1974, for example, rain falling at Pitlochry, Scotland had a pH measured at 2.4 (Last and Nicholson 1982), and a value of 2.7 was reported from western Norway a few weeks later (Sage 1980). At Dorset, north of Toronto, Ontario, snow with a pH of 2.97 fell in the winter of 1976-77 (Howard and Perley

1991), and an extreme value of pH 1.5, some 11,000 times more acid than normal, was recorded for rain falling in West Virginia in 1979 (LaBastille 1981). Although these values are exceptional, the pattern is not. Acid deposition tends to be episodic, with a large proportion of the acidity at a particular site arriving in only a few days of heavy precipitation (Last 1989). Annual averages also mask large seasonal variations. Summer precipitation, for example, is often more acid than that in the winter, although emissions are generally less in the summer.

The quality of the rain is determined by a series of chemical processes set in motion when acidic materials are released into the atmosphere. Some of the SO2 and NOX emitted will return to the surface quite quickly, and close to their source, as dry deposition. The remainder will be carried up into the atmosphere, to be converted into sulphuric and nitric acid, which will eventually return to earth as acid rain (see Figure 4.2). The processes involved are fundamentally simple. Oxidation converts the gases into acids, in either a gas or liquid phase reaction. The latter is more effective. The conversion of SO2 into sulphuric acid in the gas phase is 16 per cent per hour in summer and 3 per cent per hour in winter. Equivalent conversion rates in the liquid phase are 100 per cent per hour in summer and 20 per cent per hour in winter (Mason 1990). Despite the relatively slow conversion to acid in the gas phase, it is the main source of acid rain when clouds and rain are absent, or when humidity is low.

The rate at which the chemical reactions take place will also depend upon such variables as the concentration of heavy metals in the airborne particulate matter, the presence of ammonia and the intensity of sunlight. Airborne particles of manganese and iron, for example, act as catalysts to speed up the conversion of SO2 to sulphuric acid and sulphates. Natural ammonia may have similar effects (Ontario: Ministry of the Environment 1980). Sunshine provides the energy for the production of photo-oxidants—such as ozone (O3), hydrogen peroxide (H2O2) and the hydroxyl radical (OH)—from other pollutants in the atmosphere, and these oxygen-rich compounds facilitate the oxidation of SO2 and the NOX to sulphuric and nitric acid respectively (Cocks and Kallend 1988). The role of the photochemical component in the conversion process may account for the greater acidity of summer rainfall in many areas (Mason 1990). In the presence of water, these acids, and the other chemicals in the atmosphere, will dissociate into positively or negatively charged particles called ions. For example, sulphuric acid in solution is a mixture of positively charged hydrogen ions (cations) and negatively charged sulphate ions (anions). It is these solutions, or 'cocktails of ions', as Park (1987) calls them, that constitute acid rain.

Whatever the complexities involved in the formation of acid rain, the time scale is crucial. The longer the original emissions remain in the atmosphere, the more likely it is that the reactions will be completed, and the sulphuric and nitric acids produced. Long Range Transportation of Atmospheric Pollution (LRTAP)—transportation in excess of 500 km—is one of the mechanisms by which this is accomplished.

Air pollution remained mainly a local problem in the past. The effects were greatest in the immediate vicinity of the sources, and much of the effort of environmental groups in the 1960s and 1970s was expended in attempts to change that situation. Unfortunately, some of the changes inadvertently contributed to the problem of acid rain. One such was the tall stacks policy. In an attempt to achieve the reduction in ground level pollution required by the Clean Air Acts, the CEGB in Britain erected 200 m high smokestacks at its generating stations (Pearce 1982d). Industrial plants and power stations in the United States took a similar approach, increasing the heights of their stacks until, by 1977 more than 160 were over 150 m high (Howard and Perley 1991) and by 1981, at least 20 were more than 300 m high (LaBastille 1981). The International Nickle Company (INCO) added a 400 m superstack to its nickel smelter complex at Sudbury, Ontario in 1972 (Sage 1980). The introduction of these taller smokestacks on smelters and thermal electric power stations, along with the higher exit velocities of the emissions, allowed the pollutants to be pushed higher into the atmosphere. This effectively reduced local pollution concentrations, but caused the pollutants to remain in the atmosphere for longer periods of time, thus increasing the probability that the acid conversion processes would be completed. The release of pollutants at greater altitudes also placed them outside the boundary layer circulation and into the larger scale atmospheric circulation system with its potential for much greater dispersal through the mechanisms of LRTAP. The net result was a significant increase in the geographical extent of the problem of acid rain.

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  • Anssi
    How acid rain is formed in nature?
    7 years ago
  • Efrem Sayid
    What level sulfuric acid is on the pH scale?
    7 years ago

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