Sulfur is an important, relatively abundant, essential element. As is true of many elements, it takes part in a biogeochemical cycle discussed shortly. It is a major component of air pollution, particularly in industrialized areas, although natural sources of sulfur also contribute. Several oxidation states are encountered in environmental systems; the most stable under aerobic conditions is S(VI) as in SO3 and sulfates. The reduced form S(-II) is encountered in organic sulfides, including some amino acids, in H2S, and in metallic sulfides. It is a reduction product of sulfates under anaerobic conditions. Oxidation of sulfides produces chiefly SO2 [S(IV)] as the immediate product. Sulfur dioxide and sulfites, the salts produced when SO2 reacts with base, are reducing agents and are used as antioxidants in some foods (e.g., cut fresh fruits and vegetables, some shellfish, wine). While sulfite is generally not considered harmful, it causes asthmatic reactions, sometimes severe, in individuals who are sensitive to it. Elemental sulfur [S(0)] also occurs in nature, as well as some intermediate forms such as S2".
Figure 10-8 shows the sulfur species that are in equilibrium in solution at pH 10 as a function of pE. At this pH, S(II) (as HS") predominates below a pE of —7, and SO2" above; the crossover shifts to larger pE values as the pH decreases; in more acidic media, elemental sulfur can be a stable intermediate.
The most important interrelationships of these states of sulfur are shown in the sulfur cycle (Figure 10-9). A large amount of sulfur is released to the atmosphere as SO2 produced from the combustion of sulfur-containing fuels. The sulfur may be present in the fuel as organosulfur compounds, or as inorganic sulfide contaminants such as FeS2. In coal the amounts of each are comparable. More than twice as much SO2 has been produced from coal combustion as from the burning of petroleum in the United States in recent years; coal and petroleum together make up the source of close to 90% of the total SO2 emissions, with ore smelting being third. The contribution from coal can be expected to increase if the use of high sulfur coals becomes necessary to replace low-sulfur petroleum, especially in power generation. Other sources of sulfur oxides are volcanic activity, chemical processing, and incineration. Biomass burning can contribute small amounts of OCS, H2S, and other sulfur compounds in addition to SO2. A small amount (about 5%) of the sulfur
FIGURE 10-8 Concentration of sulfur species in solution as a function of pE at pH 10.
FIGURE 10-8 Concentration of sulfur species in solution as a function of pE at pH 10.
Insoluble metal sulfides
Insoluble metal sulfides
oxides produced in combustion processes is emitted as SO3 rather than SO2. Furthermore, catalytic exhaust converters on automobiles convert much of the sulfur emissions from automobile exhausts to SO3. Sulfur dioxide is also formed from H2S, which is released to the atmosphere from volcanoes and from anaerobic decomposition of organic matter in soils and sediments.
Organic sulfur compounds released to the atmosphere through biological processes include H2S, CS2, OCS, CH3SCH3, CH3SH, and CH3SSCH3. The first three are emitted chiefly from anaerobic water systems, while the fourth, dimethyl sulfide, is the product of ocean phytoplankton, and is the largest biogenic source of sulfur compounds in the atmosphere. Table 10-5 gives some estimates of the amount of sulfur emitted to the atmosphere per year; many of these numbers are highly approximate and, in the case of volcanic activity, sporadic.
The residence times of gaseous sulfur species in the atmosphere are generally short: on the order of 2 days for SO2 and only a little longer for SO3, which depends on washout by precipitation. The organic compounds undergo oxida-tive reactions with oxygen atoms, ozone, and free radicals such as the hydroxyl radical, and contribute to the SO2 content; their lifetimes also are no more than a few days. However, dimethyl sulfide is converted to methane sulfonic acid, CH3SO3H, which is a component of the particulate material that makes up cloud condensation nuclei and so plays a role in climate control. Another exception is COS, which is resistant to further oxidation and has a residence
Estimated Annual Sulfur Emissions'2'^
Source Sulfur emissions (g x 1012/yr)
Combustion, smelting 113
Chemical industry 29
Volcanic 28 Biogenetic
aDust and sea-spray sources not included. Estimates differing from those in the cited source can be found in G. Brasseur, J. J. Qrlando, and G. S. Tyndall, eds., Atmospheric Chemistry and Global Change, Oxford University Press, Oxford, U.K., 1999, Chapter 5.
Source: Data from R. J. Charleson, T. L. Anderson, and R. E. McDuff, in Global Biogeochemical Cycles, S. S. Butcher, R. J. Charlson, G. H. Orians, and G. V. Wolfe, eds., Academic Press, San Diego, CA, 1992.
time of about a year. Because of this longer life, COS is more uniformly distributed in the atmosphere than the other sulfur compounds and can reach the stratosphere, where it is photolyzed and oxidized to S02 and sulfate, making up the major source of stratospheric sulfate particles except on the rare occasions when volcanic activity injects products to these altitudes. Some hydrogen sulfide generated in soils and sediments undergoes reoxidation reactions to S, S02, or SO^, plus precipitation reactions, to form insoluble metal sulfides. Many of the processes are biochemical. Little H2S or other reduced sulfur is evolved from localized sources.
Sulfur dioxide is thermodynamically unstable with respect to the higher oxide S03 under natural conditions. However, in the atmosphere the reaction of S02 with 02 is relatively slow. The rate is influenced by photochemical processes and by catalysts. The most important oxidation process involves catalysis by metal salts present in water droplets or on dust particles. In fog or cloud, S02 reacts with water to form sulfurous acid H2S03. This is followed by oxidation:
In the presence of the hydroxyl radical in polluted air (see Section 5.3), the bisulfate radical is formed and then reacts with water and oxygen to form sulfuric acid and the hydroperoxyl radical.
Sulfur trioxide itself is extremely hygroscopic and immediately reacts with water vapor to form sulfuric acid; water droplets in air containing sulfur oxides are in fact a dilute solution of H2S04. This acid may be neutralized to sulfate salts by basic substances such as ammonia that may be present from industrial or natural biological processes.
Sulfuric acid is responsible for much of the corrosiveness of air in industrial localities. The effects can be noticed in the deterioration of some construction materials (H2S04 on carbonate causes decomposition to C02, Section 12.4.1) and of metals, paint, and so on.
There is good evidence that atmospheric sulfur compounds provide a cooling influence on climate. Incident solar radiation is reflected directly by sulfate salt aerosols, and cloud formation is encouraged by the additional nucleation sites that can be provided by high atmospheric sulfur levels. It appears that increased atmospheric pollution over the last few decades may have reduced the global warming effects of the greenhouse gases by a significant amount. Natural volcanic as well as anthropogenic sources contribute to this. Atmospheric sulfates are washed out with rain, but very violent volcanic release can result in their being transported to stratospheric levels, where they can remain for periods of years and participate in the chemical reactions of the upper atmosphere.
Rainwater containing H2S04 is acidic and contributes to the acidity of lakes and streams. As discussed in Section 11.4, rainwater in the northeastern United
States has shown pH values near 4 in recent years. This is well below the value expected for water in equilibrium with atmospheric C02. Not all this acidity is necessarily caused by sulfur oxides; nitrogen oxides also play a part. In addition, the pH of rainwater depends not only on the acidic oxides, but also on basic materials emitted to the atmosphere from the same or independent sources. Thus, pH may be influenced by such factors as the nature of the fuels used for heating or power generation (e.g., high or low sulfur content), the combustion temperatures in widespread use (this can influence the formation of nitrogen oxides, see Section 10.3), the nature of particulate emissions, and extensive industrial or natural emissions of other kinds.
Because emission of S02 presents a major pollution problem, considerable attention has been given to efforts to reduce these emissions. Worldwide combustion of fossil fuels produces the order of 108 tons of S02 annually, and increasing use of high-sulfur fuels requires some means of reducing the amount of S02 evolved to control pollution problems. Pretreatment of fuels represents one approach. For example, coal can be separated from its FeS2 component on the basis of different densities. Removal of organic sulfur from coal or oil requires more elaborate and costly chemical treatment. For example, the crushed coal may be extracted with a heated sodium hydroxide solution under moderate pressure to remove nearly all of the inorganic sulfur, a large portion of the organic sulfur, and many of the trace metal components.
Removal of S02 from exhaust gases is another approach to the pollution problem, and a variety of techniques has been proposed. Since S02 is an acid, it can be removed by reaction with a base such as calcium carbonate.
The calcium sulfite can be converted to sulfate (gypsum)
which has some commercial value and at any rate is easier to handle. The base most likely used in such processes is calcium carbonate or calcium magnesium carbonate, both of which available cheaply as limestone and dolomite, respectively. In one approach, the crushed carbonate is added with the fuel and the sulfate salts are precipitated by electrostatic precipitators, or washed from the gas stream by a scrubber. Alternatively, the stack gases may be scrubbed by a slurry of CaC03. This method can be very effective, but it produces large amounts of sludge containing CaS04, CaS03, unreacted carbonate, and ash, which may contain heavy metals. Disposal of this sludge raises problems, since landfill is its most likely fate.
0ther bases can be used to scrub the gases. 0ne alternative process uses Mg0 as the base because MgS03 can be decomposed at reasonably low temperatures to produce S02 as a feedstock for sulfuric acid production, with the Mg0 regenerated for reuse. Sodium hydroxide or sodium carbonate scrubbing will produce sodium sulfite and bisulfite, as well as some sulfate. These materials remain in solution, but eventually have to be disposed of. Catalytic oxidation, using a solid catalyst bed after removal of particulates, will serve to convert the SO2 to SO3 that is dissolved in water to produce a solution of sulfuric acid with commercial value. Other approaches that can give useful products to partially offset costs have been proposed but are not in general use. Although gypsum, sulfuric acid, or sulfur itself, to which the products of these adsorption processes can be converted, have large-scale use, they are available very cheaply from other sources and do not have much economic value as by-products. Calcium carbonate adsorption is the predominant approach at this time.
To be practical, such processes must operate in very large scale systems and must do so reliably with a minimum of maintenance. The volumes of gas to be handled are enormous—for example, on the order of 108 ft3/h from a 750-MW coal-fired generating plant. Many by-products such as CaSO4 have little value, and indeed their disposal may add considerable expense. The capital equipment requirements are likely to be large; economics is a major consideration in applying such treatments. Despite these problems, stack gas scrubbers seem to be the most practical means of reducing power plant SO2 emissions at present.
Dry adsorption processes for both SO2 and NO* in stack gases are available and have been applied to some facilities. Adsorption on activated coke at 100-200°C produces adsorbed H2SO4. Injection of ammonia results in catalytic decomposition of the NO* to N2 and water as well as neutralization of the sulfuric acid. The coke is then heated to 500°C, which decomposes the ammonium sulfate back to nitrogen, water, and sulfur dioxide (which can be converted to useful chemicals) and regenerates the absorber.
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