11 Control of sulfur oxide The control of particulates and VOCs is mostly accomplished by physical processes(cyclones, ESPs, filters, leakage control, vapor capture, condensation)that do not involve changing the chemical nature of the pollutant. Some particles and vOCs are chemically changed into harmless materials by combustion. This chapter and the next concern pollutants--sulfur oxides and nitroger oxides that cannot be economically collected by physical means nor rendered harmless by combustion. Their control is largely chemical rather than physical. For this reason, these two chapters are more chemically oriented than the rest of the book. Sulfur and nitrogen oxides are ubiquitous pollutants, which have many sources. SOz, SO, and NO are strong respiratory irritants that can cause health damage at high concentrations. We have NAAQS for SO and NO2. The states are required to prepare SIPs for the control of NOz and so. These gases also form secondary particles in the atmosphere, contributing to our PMio and PM25 problems and impairing visibility. They are the principal causes of acid rain. The Clean Air Act of 1990, Section 401-Acid Deposition Control, requires substantial reductions in our national emissions of both sulfur and nitrogen oxides over the next few decades 11.1 The Elementary Oxidation-Reduction Chemistry Of Sulfur And Nitrogen are significantly different, but their chemistry is quite similar, as this short section showsmethods This chapter concerns sulfur oxides; the next, nitrogen oxides. Their sources and control Both sulfur and nitrogen in the elemental state are relatively inert and harmless to humans. Both are needed for life, all animals require some N and s in their bodies. However, the oxides of sulfur and nitrogen are widely recognized air pollutants. The reduced products also are, in some cases, In parallel form, the oxidation and reduction products of nitrogen and sulfur. Reduction means the addition of hydrogen or the removal of oxygen. If we reduce nitrogen, we produce ammonia (which logically should be called hydrogen nitride, but because it had a common name before modem chemical naming systems were devised, it goes by its common name, ammonia). Similarly if we reduce sulfur, we produce hydrogen sulfide. Both hydrogen sulfide and ammonia are very strong-smelling substances, gaseous at room temperature (-60'C and. C boiling points, respectively ) and toxic in high concentrations.(High concentrations due to accidental releases often cause fatalities. These occur in the production and use of ammonia as a fertilizer and refrigerant and in the production and processing of"sour"gas and oil, which contain hydrogen ulfide. Neither ammonia nor hydrogen sulfide has been shown to be toxic in the low concentrations that normally exist in the atmosphere hen nitrogen is oxidized, nitric oxide(NO)and then nitrogen dioxide(noz) form; likewise sulfur forms sulfur dioxide(SO2)and then sulfur trioxide (so3). These are all gases at room temperature or slightly above room temperature( boiling points 21C, 34. C,-10'C, and 45C respectively ) The oxides have higher boiling points than the hydrides. Both nitrogen and sulfur can also form other oxides, but these are the ones of principal air pollution interest In the atmosphere NOz and SO3 react with water to form nitric and sulfuric acids, which then react with ammonia or any other available cation to form particles of ammonium nitrate or sulfate or some other nitrate or sulfate. These particles, generally in the 0. 1 to l-H size range, are very efficient light-scatterers they persist in the atmosphere until coagulation and precipitation remove them. They are significant contributors to urban PM1O and PM2.5 problems. They are the principal causes of acid deposition and of visibility impairment in our national parks. NO and NO2 also play a significant role in the formation of o The estimated concentrations of these materials in unpolluted parts of the worlds atmosphere are SO2, 0.2 ppb; NH3, 10 ppb; NOz, I ppb 11.2 An Overview of the sulfur problem l1-1
11－1 11 Control of Sulfur Oxides The control of particulates and VOCs is mostly accomplished by physical processes (cyclones, ESPs, filters, leakage control, vapor capture, condensation) that do not involve changing the chemical nature of the pollutant. Some particles and VOCs are chemically changed into harmless materials by combustion. This chapter and the next concern pollutants--sulfur oxides and nitrogen oxides that cannot be economically collected by physical means nor rendered harmless by combustion. Their control is largely chemical rather than physical. For this reason, these two chapters are more chemically oriented than the rest of the book. Sulfur and nitrogen oxides are ubiquitous pollutants, which have many sources. SO2, SO3, and NO2 are strong respiratory irritants that can cause health damage at high concentrations. We have NAAQS for SO2 and NO2. The states are required to prepare SIPs for the control of NO2 and SO2. These gases also form secondary particles in the atmosphere, contributing to our PM10 and PM2.5 problems and impairing visibility. They are the principal causes of acid rain. The Clean Air Act of 1990, Section 401--Acid Deposition Control, requires substantial reductions in our national emissions of both sulfur and nitrogen oxides over the next few decades. 11.1 The Elementary Oxidation-Reduction Chemistry Of Sulfur And Nitrogen This chapter concerns sulfur oxides; the next, nitrogen oxides. Their sources and control methods are significantly different, but their chemistry is quite similar, as this short section shows. Both sulfur and nitrogen in the elemental state are relatively inert and harmless to humans. Both are needed for life; all animals require some N and S in their bodies. However, the oxides of sulfur and nitrogen are widely recognized air pollutants. The reduced products also are, in some cases, air pollutants. In parallel form, the oxidation and reduction products of nitrogen and sulfur. Reduction means the addition of hydrogen or the removal of oxygen. If we reduce nitrogen, we produce ammonia (which logically should be called hydrogen nitride; but because it had a common name before modem chemical naming systems were devised, it goes by its common name, ammonia). Similarly, if we reduce sulfur, we produce hydrogen sulfide. Both hydrogen sulfide and ammonia are very strong-smelling substances, gaseous at room temperature (-60。C and -33。C boiling points, respectively), and toxic in high concentrations. (High concentrations due to accidental releases often cause fatalities. These occur in the production and use of ammonia as a fertilizer and refrigerant and in the production and processing of "sour" gas and oil, which contain hydrogen sulfide.) Neither ammonia nor hydrogen sulfide has been shown to be toxic in the low concentrations that normally exist in the atmosphere. When nitrogen is oxidized, nitric oxide (NO) and then nitrogen dioxide (NO2) form; likewise, sulfur forms sulfur dioxide (SO2) and then sulfur trioxide (SO3). These are all gases at room temperature or slightly above room temperature(boiling points 21。C, 34。C, -10。C, and 45。C, respectively). The oxides have higher boiling points than the hydrides. Both nitrogen and sulfur can also form other oxides, but these are the ones of principal air pollution interest. In the atmosphere NO2 and SO3 react with water to form nitric and sulfuric acids, which then react with ammonia or any other available cation to form particles of ammonium nitrate or sulfate or some other nitrate or sulfate. These particles, generally in the 0.1 to 1-μ size range, are very efficient light-scatterers; they persist in the atmosphere until coagulation and precipitation remove them. They are significant contributors to urban PMl0 and PM2.5 problems. They are the principal causes of acid deposition and of visibility impairment in our national parks. NO and NO2 also play a significant role in the formation of O3. The estimated concentrations of these materials in unpolluted parts of the world's atmosphere are SO2, 0.2 ppb; NH3, 10 ppb; NO2, 1 ppb. 11.2 An Overview of the Sulfur Problem
Figure 11. 1 on page 398 shows in part how sulfur moves in the environment as a result of human activities. It sulfur emitted by erupt the movement 名 of sulfur int then back out of decaying plants. Sulfur element in the arth' s crust WI abundance of about ppm. The vast ajority form 道喜 sulfates 日 mostly gypsum the principal 夏了 anhydrite CaSO4 G Inert, nontoxIc, water-solu m fow of sulfur in the envirunment, as influenced by humans. This figure omits the large amounts of mineral uble 11.1 some volcanic eruptions release, and the flow into and out of the atmosphere to growing and found widely throughout the world All organic fuels used by humans(oil, coal, natural gas, peat, wood, other organic matter)contain some sulfur. Fuels like wood have very little(0. I per center less), whereas most coals have 0.5 percent to 3 percent(see Appendix C). Oils generally have more sulfur than wood but less than coal. If we burn the fuels, the contained sulfur will mostly form sulfur dioxide S+O2-->SO2 (in fuel) If we put this into the atmosphere, it will eventually fall with precipitation, mostly in the ocean (because most of the worlds rain falls on the ocean), and over time become part of the land mass as a result of geologic processes. Again over geologic time, it will enter into fossil fuels and l1-2
11－2 Figure 11.1 on page 398 shows in part how sulfur moves in the environment as a result of human activities. It does not include the large amounts of sulfur emitted by volcanic eruptions nor the movement of sulfur into growing plants and then back out of decaying plants. Sulfur is the sixteenth-mos t abundant element in the earth's crust, with an abundance of about 260 ppm. The vast majority of this sulfur exists in the form of sulfates, mostly as gypsum, CaSO4 .2H20, the principal ingredient of plaster and wallboards, or anhydrite, CaSO4. Gypsum is a chemically inert, nontoxic, slightly water-soluble mineral, found widely throughout the world. All organic fuels used by humans (oil, coal, natural gas, peat, wood, other organic matter) contain some sulfur. Fuels like wood have very little (0.1 per center less), whereas most coals have 0.5 percent to 3 percent (see Appendix C). Oils generally have more sulfur than wood but less than coal. If we burn the fuels, the contained sulfur will mostly form sulfur dioxide, S +O2 ---> SO2 (11.1) (in fuel) If we put this into the atmosphere, it will eventually fall with precipitation, mostly in the ocean (because most of the world's rain falls on the ocean), and over time become part of the land mass as a result of geologic processes. Again over geologic time, it will enter into fossil fuels and
ulfide minerals, which humans extract and use. These uses generally lead to the formation of so If we wish to prevent this SOz from getting into the atmosphere, we can use any of the methods described in this chapter, all of which have the effect of capturing the sulfur dioxide in the form of CasO4 2H20 that will then be returned to the earth, normally in a landfill. Most often the overall reaction will be CaCO3+So2+0.502--->CaSo4+C02 (112) (limestone) In this reaction one kind of widely available rock(limestone) is mined and used to produce another rock(anhydrite or, with 2H20, gypsum), which we put back into the ground, and to release carbon dioxide to the atmosphere We are concerned about adding to the CO2 in the atmosphere but not nearly as much as we are about adding an equivalent amount of SOz. Although Eq (11.2) ppears simple, the details of carrying it out on a large scale are complex, as discussed in this chapter In natural gas most of the sulfur is in the form of H2 S, which is easily separated from the other constituents of the gas. In oil (liquid petroleum)and also in oil shales and tar sands, the sulfur is chemically combined with the hydrocarbon compounds; normally it cannot be removed without breaking chemical bonds. In oils the sulfur is concentrated in the higher-boil ing fraction of the oil, so the same crude oil can yield a low-sulfur gasoline(average 0.03%S)and a high-sulfur heavy fuel oil(e.g, 0.5 percent to I percent S). In coal much of the sulfur is also in the form of chemically bound sulfur, but some coals have a large fraction of their sulfur in the form of small (typically 100 u) crystals of iron pyrite("fools gold, "FeS2). When the fuel is burned, almost all of the sulfur in the fuel, whether chemically bound or pyritic, is converted to sulfur dioxide(so) and carried along with stack gas. Some small fraction is captured in the ash, and some is converted to SO3. Mixtures of So and SO are sometimes called SOx to remind us that some of the sulfur is in the form of SO. Usually the SO3 is negligible, and we speak of these streams as if the only ulfur oxide they contained was so2 The other important source of soz attributable to humans is the processing of sulfur-bearing ores The principal copper ore of the world is chalcopyrite, CuFeS2. The basic scheme for obtaining copper from it is the overall high-temperature smelting reaction, CuFeS2+5/2 0, Cu+FeO+2SO in which the iron is converted to a molten oxide that will float on the molten copper( with a silica flux) and thus be separated from it. The sulfur is converted to gaseous SO2. The principal ores of lead, zinc, and nickel are also sulfides, whose processing is similar to Eq(11. 3) Because the SOz liberated in the preceding process has been widely recognized as an air pollutant for many years, considerable effort has been devoted to finding other ways to process these ores that do not produce SO2. There has been some success in developing processes that treat these ores by aqueous chemistry without producing any SO2 at all. Currently such processes are economical for partly oxidized copper oxide ores containing smaller amounts of sulfur. However, for ores like chalcopyrite, the processes have not proven economical and most of these ores are currently smelted with air or oxygen The sulfur-containing gas streams most often dealt with in industry belong to three categories--reduced sulfur, concentrated SO2 streams, and dilute SO2 streams --each with its own control method. as discussed in this chapter 11.3 The Removal of Reduced Sulfur Compounds from petroleum and Natural Gas As discussed, we can convert sulfur in organic compounds to various forms by oxidation or reduction. Here we discuss the technology for removing sulfur from gas streams when the sulfur is present in reduced form. These gas streams occur in many natural gas deposits and in many by-product gases produced in oil refining and in the fuel gases produced by coal gasification This is a large liquid flow rate. To make the system practical, one must find a solvent that can absorb much more H2s than can the water. Fortunately, for many of the gases of air pollution and industrial interest, we can do that. H2S, SO, SO, NO2, HCl, and COz are acid gases, which form acids by dissolving in water. For H2S the process H2S (gas)+H2S( dissolved in water)+*H*+ HS If we can add something to the scrubbing solution that will consume either the h or the HS, then more H,s can dissolve in the water. and much less water is needed. For acid gase s, the obvious l1-3
11－3 sulfide minerals, which humans extract and use. These uses generally lead to the formation of SO2. If we wish to prevent this SO2 from getting into the atmosphere, we can use any of the methods described in this chapter, all of which have the effect of capturing the sulfur dioxide in the form of CaSO4 . 2H20 that will then be returned to the earth, normally in a landfill. Most often the overall reaction will be CaCO3 + SO2 + 0.502 ----> CaSO4 + CO2 (11.2) (limestone) In this reaction one kind of widely available rock (limestone) is mined and used to produce another rock (anhydrite or, with 2H20, gypsum), which we put back into the ground, and to release carbon dioxide to the atmosphere. We are concerned about adding to the CO2 in the atmosphere, but not nearly as much as we are about adding an equivalent amount of SO2. Although Eq. (11.2) appears simple, the details of carrying it out on a large scale are complex, as discussed in this chapter. In natural gas most of the sulfur is in the form of H2 S, which is easily separated from the other constituents of the gas. In oil (liquid petroleum) and also in oil shales and tar sands, the sulfur is chemically combined with the hydrocarbon compounds; normally it cannot be removed without breaking chemical bonds. In oils the sulfur is concentrated in the higher-boiling fraction of the oil, so the same crude oil can yield a low-sulfur gasoline (average 0.03% S) and a high-sulfur heavy fuel oil (e.g., 0.5 percent to 1 percent S). In coal much of the sulfur is also in the form of chemically bound sulfur, but some coals have a large fraction of their sulfur in the form of small (typically 100 μ) crystals of iron pyrite ("fools gold," FeS2). When the fuel is burned, almost all of the sulfur in the fuel, whether chemically bound or pyritic, is converted to sulfur dioxide (SO2) and carried along with stack gas. Some small fraction is captured in the ash, and some is converted to SO3. Mixtures of SO2 and SO3 are sometimes called SOx to remind us that some of the sulfur is in the form of SO3. Usually the SO3 is negligible, and we speak of these streams as if the only sulfur oxide they contained was SO2. The other important source of SO2 attributable to humans is the processing of sulfur-bearing ores. The principal copper ore of the world is chalcopyrite, CuFeS2. The basic scheme for obtaining copper from it is the overall high-temperature smelting reaction, CuFeS2 + 5/2 O2 → Cu + FeO + 2SO2 (11.3) in which the iron is converted to a molten oxide that will float on the molten copper (with a silica flux) and thus be separated from it. The sulfur is converted to gaseous SO2. The principal ores of lead, zinc, and nickel are also sulfides, whose processing is similar to Eq. (11.3). Because the SO2 liberated in the preceding process has been widely recognized as an air pollutant for many years, considerable effort has been devoted to finding other ways to process these ores that do not produce SO2. There has been some success in developing processes that treat these ores by aqueous chemistry without producing any SO2 at all. Currently such processes are economical for partly oxidized copper oxide ores containing smaller amounts of sulfur. However, for ores like chalcopyrite, the processes have not proven economical and most of these ores are currently smelted with air or oxygen. The sulfur-containing gas streams most often dealt with in industry belong to three categories--reduced sulfur, concentrated SO2 streams, and dilute SO2 streams --each with its own control method, as discussed in this chapter. 11.3 The Removal of Reduced Sulfur Compounds from Petroleum and Natural Gas Streams As discussed, we can convert sulfur in organic compounds to various forms by oxidation or reduction. Here we discuss the technology for removing sulfur from gas streams when the sulfur is present in reduced form. These gas streams occur in many natural gas deposits and in many by-product gases produced in oil refining and in the fuel gases produced by coal gasification. This is a large liquid flow rate. To make the system practical, one must find a solvent that can absorb much more H2S than can the water. Fortunately, for many of the gases of air pollution and industrial interest, we can do that. H2S, SO2, SO3, NO2, HCI, and CO2 are acid gases, which form acids by dissolving in water. For H2S the process is H2S (gas) ↔ H2S (dissolved in water) ↔ H+ + HS- (11.4) If we can add something to the scrubbing solution that will consume either the H+ or the HS- , then more H2S can dissolve in the water, and much less water is needed. For acid gases, the obvious
choice is some alkali, a source of OH that can remove the h by H++OH←H20 (11.5) Removing the H+ on the fight side of Eq.(11. 4)drives the equilibrium to the fight, greatly increasing the amount of H,S absorbed The Uses and Limitations of Absorbers and Strippers for Air Pollution Control Absorber-stripper combinations are widely used to remove HCs from exhaust gas streams. This example shows that the removal of H2S from natural gas and similar streams is simple and straightforward. The system also works extremely well for removing ammonia from a gas stream, because NHs is very soluble in water or in weak acids, forming a weak alkali by the following reaction NH3+H20→NH4+OH (11.7) It is possible to make practically complete removal of NH3 from gas streams with water or weak acids. The solubility of ammonia is so high that generally the simplest possible forms of this arrangement are satisfactory. To remove SOz from gas streams by this method is also relatively easy if there are no other acid gases present. For example, SO2 could be easily removed from N2 by the scheme using any weak alkali(for example, ammonium hydroxide), and the solution would be easily regenerated to produce pure SO2. The problem of removing sulfur dioxide from combustion gases is much more complex and difficult, as discussed NO and NOz are not readily removed from gas streams by the process. Although NOz is an acid gas that produces nitric acid by reaction with water 3NO2+H20÷2HNO3+NO (11.8) the reaction rate is slow. NO is not an acid gas, so that although we can remove NO2 from a gas stream with an alkaline solvent, we cannot remove No with the same solvent. For this reason, weak alkali solvents are not successful for the joint removal of NO and NO2 or for the rapid removal of NO2 alone. No other solvent is known that serves well for this task. (My generation has not found a suitable solvent to do this; fame and fortune await the person who finds a suitable solvent to remove NO, NO2, and SOz economically from combustion gases scheme is widely used in the chemical and petroleum industries to make separations not directly related to pollution control, e.g, the separation of COz from H. The absorption column can also be used without regenerating the absorbent solution if the amount of material to be collected is small and there is some acceptable way of disposing of the loaded absorbent Sulfur Removal from Hydrocarbons Once H2S has been separated from the other components of the gas, it is normally reacted with oxygen from the air in controlled amounts to oxidize it only as far as elemental sulfur, H2S+1/202→S+H2O and not as far as so H2S+3/202→SO2+H2O (11.10) The elemental sulfur is either sold for use in the production of sulfuric acid or land-filled if them is no nearby market for it Although the chemical reaction in Eq. (11. 9) for production of sulfur(the Claus process) is simple enough, there are a variety of Ways of carrying it out, and the details can be complex; see Kohl and Nielsen. Hundreds of such plants operate successfully throughout the world; every major petroleum refinery has at least one Because elemental sulfur is inert and harmless and because reduced sulfur in the form of hydrogen ulfide or related compounds can be easily oxidized to sulfur or sulfur oxides, the entire strategy of the petroleum and natural gas industries in dealing with reduced sulfur in petroleum, natural gas, and other process gases is to keep the sulfur in the form of elemental sulfur or reduced sulfur(for example, H2S). Oxygen from the air is virtually free, so we can always move in the oxidation direction at low cost. In contrast, hydrogen is an expensive raw material, so that moving in the reduction direction is expensive Sulfur in hydrocarbon fuels(natural gas, propane, gasoline, jet fuel, diesel fuel, furnace oil)is normally converted to So during combustion and then emitted to the atmosphere. Larg oil-burning facilities can have equipment to capture that SOz, but autos, trucks, and airplanes do not. The only way to limit the SO2 emissions from these sources is to limit the amount of sulfur in the fuel. For this reason the Clean air Act of 1990 lim its the amount of sulfur in diesel fuel to 0.05 percent by weight. Crude oils vary in their sulfur contents: low-sulfur crudes are called"sweet high-sulfur crudes, "sour. "If the fraction of the crude oil going to gasoline or diesel fuel has too l1-4
11－4 choice is some alkali, a source of OHthat can remove the H+ by H+ +OH- ↔ H20 (11.5) Removing the H+ on the fight side of Eq. (11.4) drives the equilibrium to the fight, greatly increasing the amount of H2S absorbed. The Uses and Limitations of Absorbers and Strippers for Air Pollution Control Absorber-stripper combinations are widely used to remove HCs from exhaust gas streams. This example shows that the removal of H2S from natural gas and similar streams is simple and straightforward. The system also works extremely well for removing ammonia from a gas stream, because NH3 is very soluble in water or in weak acids, forming a weak alkali by the following reaction: NH3 + H20 → NH4 + + OH- (11.7) It is possible to make practically complete removal of NH3 from gas streams with water or weak acids. The solubility of ammonia is so high that generally the simplest possible forms of this arrangement are satisfactory. To remove SO2 from gas streams by this method is also relatively easy if there are no other acid gases present. For example, SO2 could be easily removed from N2 by the scheme using any weak alkali (for example, ammonium hydroxide), and the solution would be easily regenerated to produce pure SO2. The problem of removing sulfur dioxide from combustion gases is much more complex and difficult, as discussed. NO and NO2 are not readily removed from gas streams by the process. Although NO2 is an acid gas that produces nitric acid by reaction with water, 3NO2 + H20 → 2HNO3 + NO (11.8) the reaction rate is slow. NO is not an acid gas, so that although we can remove NO2 from a gas stream with an alkaline solvent, we cannot remove NO with the same solvent. For this reason, weak alkali solvents are not successful for the joint removal of NO and NO2 or for the rapid removal of NO2 alone. No other solvent is known that serves well for this task. (My generation has not found a suitable solvent to do this; fame and fortune await the person who finds a suitable solvent to remove NO, NO2, and SO2 economically from combustion gases by the scheme) The scheme is widely used in the chemical and petroleum industries to make separations not directly related to pollution control, e.g., the separation of CO2 from H2. The absorption column can also be used without regenerating the absorbent solution if the amount of material to be collected is small and there is some acceptable way of disposing of the loaded absorbent. Sulfur Removal from Hydrocarbons Once H2S has been separated from the other components of the gas, it is normally reacted with oxygen from the air in controlled amounts to oxidize it only as far as elemental sulfur, H2S+ 1/2 O2 → S + H2O (11.9) and not as far as SO2, H2S + 3/2 O2 → SO2+H2O (11.10) The elemental sulfur is either sold for use in the production of sulfuric acid or land-filled if them is no nearby market for it. Although the chemical reaction in Eq. (11.9) for production of sulfur (the Claus process) is simple enough, there are a variety of Ways of carrying it out, and the details can be complex; see Kohl and Nielsen. Hundreds of such plants operate successfully throughout the world; every major petroleum refinery has at least one. Because elemental sulfur is inert and harmless and because reduced sulfur in the form of hydrogen sulfide or related compounds can be easily oxidized to sulfur or sulfur oxides, the entire strategy of the petroleum and natural gas industries in dealing with reduced sulfur in petroleum, natural gas, and other process gases is to keep the sulfur in the form of elemental sulfur or reduced sulfur (for example, H2S). Oxygen from the air is virtually free, so we can always move in the oxidation direction at low cost. In contrast, hydrogen is an expensive raw material, so that moving in the reduction direction is expensive. Sulfur in hydrocarbon fuels (natural gas, propane, gasoline, jet fuel, diesel fuel, furnace oil)is normally converted to SO2 during combustion and then emitted to the atmosphere. Large oil-burning facilities can have equipment to capture that SO2, but autos, trucks, and airplanes do not. The only way to limit the SO2 emissions from these sources is to limit the amount of sulfur in the fuel. For this reason the Clean Air Act of 1990 limits the amount of sulfur in diesel fuel to 0.05 percent by weight. Crude oils vary in their sulfur contents: low-sulfur crudes are called "sweet"; high-sulfur crudes, "sour." If the fraction of the crude oil going to gasoline or diesel fuel has too
high a sulfur content( which many do under current regulations), most of that sulfur is removed by catalytic hydrodesulfurization, The mixture leaving the reactor is cooled, condensing most of the hydrocarbons. The remaining gas stream, a mixture of H2 and HS, is one of the streams treated in a refinery for H2S removal by the process. Some petroleum streams in refineries are treated over these catalysts to remove both sulfur and nitrogen because those elements interfere with the catalysts used for subsequent processing. The resulting gas streams contain both H2S and NH Whether the treatment of gases with high concentrations of H2S and NH3 should be considered as air pollution control is an open question. For natural gas fields with H2S, treatment is a market requirement, because the typical purchase specification for natural gas in the United States is H2S <4 ppm. However, at one time in oil refineries H2S-containing gases were customarily burned for internal heat sources in the refineries if the H2s content was modest. Current U.S. EPa air pollution regulations forbid the burning of such refinery waste gases if they contain more than 230 mg/dscm(dry standard cubic meter)of H2S, so the removal of H2s down to that concentration in oil refinery gases is done by the method to meet air pollution control regulations 11.4 Removal of SOz From Lean Waste Gases The major source of So, except near uncontrolled copper, lead, zinc, and nickel smelters, which no longer exist in the United States, but do in some developing countries, is the stacks of large coal-or oil-burning facilities. Most of the largest ones are coal-burning electric power plants. For them, the typical Soz content of the exhaust gas is about 0. 1 percent SOz, or 1000 ppm(see Example 7.10), which is much too low for profitable recovery as H2SO There are several drawbacks to this procedure for dealing with the sO from an electric power plant. First, it requires a large amount of water. The computed water flow is approximately 1 percent of the flow of the Hudson River at New York City. Power plants located on the Hudso the Mississippi, the Ohio, or the Columbia rivers could obtain such amounts of water, but most of the power plants in the world could not Second, the waste water stream, which is 80 percent saturated with SO, would emit this sOz back into the atmosphere at ground level(river level causing an SO2 problem that might be more troublesome than the emission of the same amount of sO2 from the power plant's stack. Third, in aqueous solution SO2 undergoes Reaction(11 12 (without the catalyst)which would remove most of the dissolved Oz in the river, making it impossible for fish to live in it For this reason alone, simple dissolution of large quantities of so in most rivers is prohibited However, the first large power plant to treat its stack gas for SO2 removal did remove SO with river water. The Battersea Plant of the London Power Company is located on the banks of the Thames River, which is large enough to supply the water it needed. Furthermore, the water of the Thames is naturally alkaline because its course passes through many limestone formations, so that it will absorb substantially more SO2 than would pure water. To prevent the dissolved SO2 from consuming O2 in the river, the effluent from the gas washers was held in oxidizing tanks, where air was bubbled through it until the dissolved SO2 was mostly oxidized to sulfate (SO42),before being discharged to the Thames. In this form the sulfur has a low vapor pressure and does not reenter the air nor kill the fish by consuming the river's dissolved oxygen. Although this pioneering plant had its problems, it was a technical success--removing over 90 percent of the SOz--and operated from 1933 to 1940. (The SOz removal system was shut down in 1940 because the exhaust plume from this plant was wet due to the scrubber and, hence, very visible. It made a good navigation marker for German aircraft during the Battle of Britain. As we saw in Example 11.2, the amount of scrubbing water required can be substantially reduced if we add a reagent to the water that increases the solubility of the gas being removed Comparing this problem to the H2S removal problem in Examples 11. I and 11. 2, we see that: 1. The volumetric flow rate of the gas is about 1700 times that in the H2S removal problem(14 times because of the higher molar flow rate, and 120 times because of the lower gas density At 2. The power cost to drive the gas through the scrubber is thus 1700 times as large, for an equal AP. Thus minimizing pressure drop is much more important in this problem than in that 3. Here there is no regenerator. If we regenerated the solution to produce a stream of practically pure So2 we would have no economical way of converting it to a harmless solid, as the claus process does with H2S In the previous examples we said little about the internal features of the absorbing column For the l1-5
11－5 high a sulfur content (which many do under current regulations), most of that sulfur is removed by catalytic hydrodesulfurization, The mixture leaving the reactor is cooled, condensing most of the hydrocarbons. The remaining gas stream, a mixture of H2 and H2S, is one of the streams treated in a refinery for H2S removal by the process. Some petroleum streams in refineries are treated over these catalysts to remove both sulfur and nitrogen because those elements interfere with the catalysts used for subsequent processing. The resulting gas streams contain both H2S and NH3. Whether the treatment of gases with high concentrations of H2S and NH3 should be considered as air pollution control is an open question. For natural gas fields with H2S, treatment is a market requirement, because the typical purchase specification for natural gas in the United States is H2S ≤ 4 ppm. However, at one time in oil refineries H2S-containing gases were customarily burned for internal heat sources in the refineries if the H2S content was modest. Current U.S. EPA air pollution regulations forbid the burning of such refinery waste gases if they contain more than 230 mg/dscm (dry standard cubic meter) of H2S, so the removal of H2S down to that concentration in oil refinery gases is done by the method to meet air pollution control regulations. 11.4 Removal of SO2 From Lean Waste Gases The major source of SO2, except near uncontrolled copper, lead, zinc, and nickel smelters, which no longer exist in the United States, but do in some developing countries, is the stacks of large coal- or oil-burning facilities. Most of the largest ones are coal-burning electric power plants. For them, the typical SO2 content of the exhaust gas is about 0.1 percent SO2, or 1000 ppm (see Example 7.10), which is much too low for profitable recovery as H2SO4. There are several drawbacks to this procedure for dealing with the SO2 from an electric power plant. First, it requires a large amount of water. The computed water flow is approximately 1 percent of the flow of the Hudson River at New York City. Power plants located on the Hudson, the Mississippi, the Ohio, or the Columbia rivers could obtain such amounts of water, but most of the power plants in the world could not. Second, the waste water stream, which is 80 percent saturated with SO2, would emit this SO2 back into the atmosphere at ground level (river level); causing an SO2 problem that might be more troublesome than the emission of the same amount of SO2 from the power plant's stack. Third, in aqueous solution SO2 undergoes Reaction (11.12), (without the catalyst) which would remove most of the dissolved 02 in the river, making it impossible for fish to live in it. For this reason alone, simple dissolution of large quantities of SO2 in most rivers is prohibited. However, the first large power plant to treat its stack gas for SO2 removal did remove SO2 with river water. The Battersea Plant of the London Power Company is located on the banks of the Thames River, which is large enough to supply the water it needed. Furthermore, the water of the Thames is naturally alkaline because its course passes through many limestone formations, so that it will absorb substantially more SO2 than would pure water. To prevent the dissolved SO2 from consuming O2 in the river, the effluent from the gas washers was held in oxidizing tanks, where air was bubbled through it until the dissolved SO2 was mostly oxidized to sulfate (SO4 2- ), before being discharged to the Thames. In this form the sulfur has a low vapor pressure and does not reenter the air nor kill the fish by consuming the river's dissolved oxygen. Although this pioneering plant had its problems, it was a technical success--removing over 90 percent of the SO2--and operated from 1933 to 1940. (The SO2 removal system was shut down in 1940 because the exhaust plume from this plant was wet due to the scrubber and, hence, very visible. It made a good navigation marker for German aircraft during the Battle of Britain.) As we saw in Example 11.2, the amount of scrubbing water required can be substantially reduced if we add a reagent to the water that increases the solubility of the gas being removed. Comparing this problem to the H2S removal problem in Examples 11.1 and 11.2, we see that: 1. The volumetric flow rate of the gas is about 1700 times that in the H2S removal problem (14 times because of the higher molar flow rate, and 120 times because of the lower gas density At 100 atm, methane has m 1.2 times the density of a perfect gas) 2. The power cost to drive the gas through the scrubber is thus 1700 times as large, for an equal AP. Thus minimizing pressure drop is much more important in this problem than in that. 3. Here there is no regenerator. If we regenerated the solution to produce a stream of practically pure SO2 we would have no economical way of converting it to a harmless solid, as the Claus process does with H2S. In the previous examples we said little about the internal features of the absorbing column. For the