Classroom Illustrations of Acidic Air Pollution Using ... - ACS Publications

Chemistry for Everyone

Classroom Illustrations of Acidic Air Pollution Using Nylon Fabric Dean J. Campbell,* Emily A. Wright, Mardhia O. Dayisi, Michael R. Hoehn, Branden F. Kennedy, and Brian M. Maxfield Department of Chemistry and Biochemistry, Bradley University, Peoria, Illinois 61625, United States *[email protected]

June 10, 1970, seemed to be a fairly typical warm and muggy summer day in central Illinois, but something out of the ordinary arrived from the skies. In downtown Peoria by the riverfront, businesses were making money, factories were making products, and smoke billowed into the sky from the smokestack of one of the coal-burning power plants. When lunchtime rolled around, the employees from the local stores walked to the local restaurants to refuel for the rest of the afternoon. According to an article in the local paper, the “Peoria Journal Star” (1), the ladies working downtown noticed ash deposits and holes in their nylon stockings after returning from lunch. They concluded that the ash had attacked the fibers of the nylon stockings. Many ladies returned their stockings to the store where they had been purchased, but the same disintegration occurred as they walked back to work (1). When the employees returned to work, many of them reported having headaches and burning of the eyes.1 The affected people were angered at the damage apparently caused by this ash, even though pollution regulations were weaker in those days. Accusations were made against the power plant, but a company representative denied the allegations by saying, “We haven't changed anything we are doing that would cause this phenomenon, but we will investigate to see if there is something that we aren't aware of that would have caused it” (1). The local news featured a company representative placing stockings under the smoke stack and no damage was observed (2). The company claimed that because of this, there was no way that the smoke could be blamed. Many people took this as fact. However, what happened next became part of the Bradley University Chemistry departmental lore. David Sweet, a student who was researching with Professor Tom Cummings in the Chemistry Department, believed that this demonstration was flawed. The ash that had fallen from the sky was fly ash, a byproduct of coal combustion flying up and out the power plant smokestack. It appeared to him that the fly ash falling from the smokestack had carried sulfur-containing oxyacids out of the exhaust plume down to earth. When the ash came in contact with the nylon stockings, the acids attacked the fabric. David Sweet was so infuriated by the power company's claims that he called both the company and a local TV station to complain. There was no success with the company, but a TV news reporter asked him for an interview. Sweet went to the lab to come up with a demonstration proving the company was wrong (2). He ran a sequence of experiments in which he passed sulfur dioxide from a small tank up a ceramic tube through nylon stockings in attempt to mimic the environmental conditions on the day the nylon damage occurred. The first experiment used a dry, room-temperature stocking, and as he passed the sulfur dioxide through it there was no damage. He then passed the sulfur dioxide

_

through a pair of stockings that were dampened by a slight aqueous mist from a spray bottle; still there was no damage. Sweet sprinkled a dry pair of stockings with transition-metal oxides, such as iron, chromium, and manganese oxides, and passed the sulfur dioxide and there was still no damage. Finally, Sweet dampened a pair of stockings, sprinkled them with transition-metal oxides, and passed the sulfur dioxide through them, and they disintegrated. This reaction with the damp stockings showed how the weather and pollutants present in the air could be in the right balance to cause the deterioration of the stockings (2). The TV news crew filmed the demonstration and presented it to the public. How the Events Happened Industrial Production of Sulfuric Acid Industrially, sulfur-containing oxyacids are produced on a massive scale. Sulfurous acid can be produced by first combining sulfur and oxygen to produce sulfur dioxide (e.g., by burning sulfur or heating sulfide ores in an excess of air): ð1Þ S þ O2 f SO2 The sulfur dioxide can then be combined with water to form sulfurous acid: ð2Þ H2 O þ SO2 f H2 SO3 Sulfuric acid is produced in a process called the contact process. Here, sulfur dioxide combines further with oxygen to form sulfur trioxide: ð3Þ 2SO2 þ O2 h 2SO3 The formation of sulfur trioxide is reversible and can be very slow. For the industrial-scale production of sulfur trioxide, vanadium(V) oxide is used to catalyze the oxidation reaction at about 450 K (3, 4). Hypothetically, sulfuric acid could form by simply reacting sulfur trioxide directly with water: H2 O þ SO3 f H2 SO4

ð4Þ

However, this is a dangerously exothermic reaction. To circumvent this, the sulfur trioxide is dissolved in concentrated sulfuric acid, which produces fuming sulfuric acid (also called oleum). The fuming sulfuric acid can be more safely reacted with water to make more quantities of concentrated sulfuric acid (4). Sulfuric Acid from Coal Combustion At coal-burning power plants, sulfur-containing compounds can be converted to sulfur-containing oxyacids by a number of routes. There are also multiple potential sources of this sulfur. For example, iron sulfides such as marcasite and pyrite

_

r 2011 American Chemical Society and Division of Chemical Education, Inc. pubs.acs.org/jchemeduc Vol. 88 No. 4 April 2011 10.1021/ed100604a Published on Web 02/14/2011

_

Journal of Chemical Education

387


carboxylic acid functional groups. There are different types of nylon polymers, but a common type used to make fibers for women's stockings is nylon-6,6, which is made by polymerizing hexanedioic acid (C6H8O4) and 1,6-diaminohexane (C6H16N2) (13). The two monomers each contain six carbon atoms, hence, the name of the product, nylon-6,6. This condensation reaction produces water as the amide linkages are formed. The reverse reaction can occur, resulting in the hydrolysis of the amide group by water molecules in the presence of acid catalysts and heat to produce the monomers.

Figure 1. Concentrated sulfuric acid attacks a swatch of nylon fabric.

(both with the formula FeS2) can occur in coal and will burn to produce iron oxides and sulfur oxides (5). The conversion of sulfur dioxide to sulfur trioxide can occur within the flames of the boilers of the power plant, but this exothermic reaction is not favored at these high temperatures (5). This reaction can also take place at somewhat lower temperatures with the catalytic assistance of iron oxides on the fly ash and on the surfaces of equipment in the plant (5, 6). Sulfur trioxide in the flue gases combines with water vapor either in the flue gases within the smokestack or with water vapor in the atmosphere outside of the smokestack. At temperatures between about 370 and 425 K, the sulfuric acid will condense in the air as droplets or on surfaces (5). Sulfur dioxide can also be converted to sulfuric acid by other routes. Sulfur dioxide that leaves the smokestack of the power plant can react with ultraviolet light in sunlight and other species in the air such as hydroxyl radicals, biacetyl, benzaldehyde, and nitrogen dioxide to produce sulfur trioxide, but whether the sulfur dioxide can be oxidized in sunlight without these species present is controversial (7, 8). The sulfur dioxide can react with liquid water adsorbed onto the ash surfaces to produce sulfurous acid. The sulfite ions in this sulfurous acid solution can then be oxidized to sulfate ions to make sulfuric acid. This oxidation process can be catalyzed by iron(III) species in the solution; the ash itself could act as a source of these ions (9). Hydrogen sulfite can also be photochemically oxidized in an aqueous solution containing solid iron(III) oxide (8). The metal oxides in the fly ash come from metal compounds in the coal that have reacted with the oxygen during the combustion process (10). Analysis of fly ash often yields combinations of eight oxide components (6) in varying concentrations: SiO2, Al2O3, Fe2O3, CaO, MgO, Na2O, K2O, and SO3. The supporting information contains a scanning electron microscope (SEM) image and electron dispersion of X-rays (EDX) analysis of a recent sample of fly ash from a coal-burning power plant. Nylon Nylon was first discovered and patented by Wallace Carothers and his research group at the DuPont Experimental Station (11, 12). Nylon is a polyamide, containing amide functional groups made by the condensation reactions of amine and 388


_

Vol. 88 No. 4 April 2011

_

Studies have been performed to analyze the effect of acid on nylon. Research showed that nylon fibers degraded similarly when exposed to various concentrations of either hydrochloric or sulfuric acid at 50 °C (14). By placing the fibers in simulated environmental conditions, researchers have shown that the most significant damage to nylon fibers occurred when they were wet and exposed to light and 0.2 ppm sulfur dioxide. Other research explored the degradation of nylon fibers in varying acid conditions (ranging from distilled water to 1.0 M sulfuric acid) and temperature conditions (20-90 °C). This research showed that an increase in temperature increases the absorption of water by the nylon fibers and consequently their acid degradation (15). Bringing the Events to the Classroom Existing Experiments Sweet's experiment involved a small tank of toxic and corrosive sulfur dioxide, and we have had difficulties directly reproducing the results from the sketchy details of this experiment. The nylon samples have not disintegrated in the short time scales that we have run the experiments, as we are unwilling to send extensive quantities of sulfur dioxide up the flues of our fume hoods. However, classroom or laboratory demonstrations can be performed to illustrate portions of the overall chemical event. A sample of FeS2 (believed to be marcasite) found near a central Illinois coal seam was ground and heated in a loosely corked test tube over a Bunsen burner. The solid decomposed, releasing sulfur that condensed on the inner walls of the tube. Upon further heating, the sulfur disappeared. Wet pH paper placed into the tube turned red, indicating the presence of acidic vapor, most likely sulfur dioxide. Other methods of sulfur dioxide production involve the combination of sulfite or hydrogen sulfite salts with acids to form sulfurous acid, which decomposes to produce sulfur dioxide (16, 17). Perhaps the simplest way to produce sulfur dioxide is by burning sulfur in air (18). A booklet that is available online describes simple environmental experiments, including one involving placing a nylon

pubs.acs.org/jchemeduc

_

r 2011 American Chemical Society and Division of Chemical Education, Inc.


stocking outdoors and inspecting it occasionally for holes that might be caused by acidic fly ash (19). To be an effective demonstration, the experiment seems to be limited to locations where acidic ash would be present in sufficient concentrations to be a nuisance. Other demonstrations of the effects of acid pollution on the environment have been published. One study looked at three major types of stone that are affected by acid rain: marble (limestone), sandstone, and granite, which are used frequently in monuments and buildings (20). Laboratory runoff experiments have been conducted to quantify the erosion of marble and limestone by acid rain (21). Aquatic life is also affected by acid rain when it interacts with Al(OH)3 found in soil and clay, causing Al3þ runoff to be introduced into bodies of water where it becomes harmful (22). Demonstration 1 A more graphic (and quite simple) demonstration involves dissolving holes in nylon stockings with drops of sulfuric acid. First, lay a small piece of nylon stocking flat in a Petri dish. Then, using an eyedropper, place a few droplets of the sulfuric acid on the stockings. The minimum concentration that seems to successfully dissolve the nylon stocking threads in a reasonable time is 4 M. The higher the concentration of acid, the faster the threads dissolve: 6 M acid works well, and concentrated sulfuric acid dissolves through the fabric, as illustrated in Figure 1. It sometimes takes a little time for the fabric to dissolve. Therefore, it is recommended that when doing the demonstration, add acid to the fabric first and then explain the connection to the nylon degradation event in Peoria while waiting for dissolution to take place. The nylon fabric can change color in the vicinity of the holes dissolved in the fabric. The very low pH of the acid droplets appears to shift the colors of the fabric dyes, much like acid-base indicators. However, the specific color the fabric turns can be unpredictable; one tan-colored stocking has turned red, and another brand has turned blue, presumably because different dyes were used to achieve the specific tan colors. This demonstration can be shown at a variety of grade levels and can be shown to large groups with an overhead projector. If the demonstration is not performed on an overhead projector, it is easier to see when the nylon fabric is placed over a color-contrasting background. Water can be used to clean up the demonstration, but some surfaces can require a bit of scrubbing to remove the sticky, gummy, partiallydissolved nylon. Demonstration 2 We have also developed a variation on Sweet's original experiment that does not require a sulfur dioxide gas tank. This open system2 still produces some sulfur dioxide and must be performed in a fume hood. To perform this experiment, shown in Figure 2, combine 7.0 g of sodium bisulfite and 200 mL of water in a 500 mL Florence flask. (We used ACS reagent grade sodium bisulfite, a mixture of NaHSO3 and Na2S2O5.) Wipe the mouth of the flask dry to remove any possible chemical contamination. Stretch about four stacked ∼3 cm squares of dry (or water-soaked3) nylon stocking fabric tightly over the mouth of the flask and then secure the fabric to the flask mouth with coated or uncoated wire. Sprinkle approximately 0.03 g of iron(III) oxide powder onto the fabric. Some powder will likely fall through the fabric layers into the sodium bisulfite solution, but this does not affect the reaction. Place the flask assembly on a hot plate to boil


_

Figure 2. Demonstration of nylon fabric degradation. The fabric is sprinkled with iron(III) oxide powder and the solution in the Florence flask contains sodium bisulfite.

the solution, releasing some warmth, water vapor, and sulfur dioxide up into the fabric at the mouth of the flask. Before the solution in the flask boils, use a clamp to hold a 400 mL beaker upside down over nylon at the mouth of the flask to help keep any vapors that pass through the nylon fabric layers in the vicinity of those layers. Tilt the beaker at an angle of roughly 20° from vertically upside down so that any condensation droplets that collect within the beaker will move along the interior beaker walls and not drip onto the nylon fabric. Rather, the droplets will move down the walls of the flask and onto the hot plate, flashing to steam. Over the time scale of a few hours (be careful not to boil the solution to dryness) the nylon fabric can be significantly damaged: often exhibiting a color change and breaking threads. Even if no damage appears within the first few hours, it may appear later if the entire experimental setup is left to cool overnight. The reason for this might be due to be a slow reaction between the acid and the nylon fibers or it might be due to a change in acid concentration on the fibers as the moisture on the fabric dries out overnight. The threads appear to break most near where the fabric meets the lip of the flask, where the threads curve the most and are under the most stress. Different samples of nylon fabric appear to have varying susceptibility to attack by the vapors, but a lack of sodium bisulfite in solution (and therefore no sulfur dioxide production) results in no nylon damage. Adding powdered iron(III) oxide to the nylon is much more damaging than adding no oxide at all. It is hypothesized that the iron(III) oxide catalyzes the formation of sulfur trioxide (and therefore sulfuric acid) at the nylon, increasing the fabric damage. Adding the iron(III) oxide directly to the NaHSO3 solution rather than the fabric did not produce degradation of the nylon. Iron(III) oxide has been the best catalyst for these experiments. Adding powdered vanadium(V) oxide, used in industrial production of


_


_


389


sulfur dioxide, does not typically produce nearly as significant damage as does powdered iron(III) oxide and is also not recommended for classroom use due its toxicity. Fly ash from a local power plant also does not typically produce nearly as significant damage as does powdered iron(III) oxide. Powdered iron(III) oxide that has been heated and then cooled does not reproducibly produce more significant damage than just the powdered oxide from the reagent bottle and often produces less damage. A sample of powdered FeS2 (believed to be marcasite) does appear to be able to produce nylon damage. This powder can also help damage nylon after it has been preheated on a hot plate and then cooled, presumably producing iron oxides, before being placed on the fabric. Hazards Safety precautions such as eye and skin protection must be observed. Sulfuric acid is very corrosive. Sodium bisulfite can produce toxic or corrosive gases upon exposure to heat or acids. Iron(II) sulfide (marcasite or pyrite) can produce toxic or corrosive gases upon exposure to heat or acids. Iron(III) oxide is fairly inert. However, the bright red powder can dust surfaces (clothing, countertops, sinks, etc.) quite effectively and can be difficult to remove. Discussion Ultimately, there were no significant repercussions or reparations resulting from David Sweet's demonstration, even though the power company was shown to be in error. The TV station footage of the demonstration has been lost to time. The reason this mysterious stocking-damaging event is rather unique in Peoria history is not well understood. Perhaps there was an extraordinary quantity of iron oxide and sulfur oxide produced in the plant emissions that day. However, a 1967 article in the “Peoria Journal Star” (23) shows the aforementioned power company was aware of emissions problems at the downtown Peoria station, which was built in 1890. The historical anecdote and accompanying demonstrations make dramatic illustrations of concepts of the acidity of nonmetal oxides, catalytic behavior, and air pollution. Air pollution will continue as nations develop industrially but what is done to reduce the pollution will influence the health, wealth, and well being of those nations. The effects of sulfur oxide pollution from power plants extend beyond a few torn nylon stockings. It affects the air we breathe, leading to health problems ranging from skin, eye, and upper respiratory irritation in low concentrations (recall the headaches and eye irritation experienced by passers-by that day), to asthma, edema of the lungs, and even respiratory paralysis in higher concentrations (24). This is why sulfur oxide-producing reactions should be handled in a fume hood. Acid precipitation increases the hydrogen ion and aluminum ion concentrations in waterways, harming life forms in the water (25). The Clean Air Act and its subsequent amendments (including one in 1970) set standards on the emission of air pollution from urban, industrial, and motor vehicle sources to protect the air quality and public health of the United States (26). Power companies have responded in a variety of ways to these standards. The age of the downtown Peoria plant and the cost of rebuilding to acceptable standards led the power company to close the station on May 2, 1971 (23, 27). For other power plants, these responses included installing scrubber systems and burning coal 390


_


_

with lower sulfur content. Though Peoria had many local coal mines in the past, some of the central Illinois coal-fired power plants prefer to burn significant quantities of lower-sulfur coal delivered by rail from the Powder River Basin in the state of Wyoming (27). One of the open-pit coal mines in the basin is actually crossed by Interstate 90 west of Gillette, WY, over 800 mi (1300 km) from Peoria. As a result of the sulfur oxide regulations, the quantity of sulfur dioxide emission in Peoria decreased from 32 ppb in 1972 to 7 ppb in 1989 (28). The average level of sulfur dioxide in Peoria county declined from 7 ppb in 1995 to 2 ppb in 2007 (29). However, complicated issues associated with sulfur oxide emissions persist. For example, oxidation of sulfur dioxide to sulfates can contribute to the growth of tiny water droplets in the atmosphere. These sulfate-containing aerosols can reflect sunlight away from the earth, producing a cooling effect (30-32). Some people have proposed deliberately adding sulfur dioxide to the atmosphere to increase the degree of sulfate aerosol cooling of the earth in an effort to combat global warming (32). Another example of the complexities of sulfur oxide emission involves fly ash, which can be removed from power plant emissions by electrostatic precipitation. This process can be facilitated by a sulfur trioxide flue gas conditioning (FGC) system where sulfur dioxide is converted to sulfur trioxide and deliberately added to the flue gases in the chimney to make them more electrically conductive in order to capture more fly ash (8, 33). Selective catalytic reduction (SCR) systems, designed to decrease nitrogen oxide emissions from power plants, have sometimes assisted the conversion of sulfur dioxide to sulfur trioxide (5, 34). In 2004, a catalyst problem of this type caused a power plant in Indiana to produce sulfur oxide emissions that blew into the community of Mt. Carmel, IL. Residents there encountered physical problems (e.g., eye irritation) similar to that encountered by the Peoria residents in 1970 (1, 34). Sulfur oxide emissions are not an issue restricted to Illinois. The global problem of the acidification of the environment from sulfur oxides is becoming more apparent as more nations such as India and China become increasingly industrialized. China currently burns more coal for energy than the United States and European Union combined and builds more coal-fired plants at the rate of about one per week (34) to meet the energy needs of its population of over one billion. There have been complaints from Japan and South Korea about increases in the concentration of sulfur dioxide in their air as a result of cross-border contamination from China (35, 36). United Sates satellites and groundbased detectors in California, Oregon, and Washington have also detected Asian pollutants wafting into North America from across the Pacific Ocean (35, 37). Sulfur oxide acidification of the environment has been and will continue to be an issue for some time. Acknowledgment We would like to thank Robert Gayhart, Max Taylor, Thomas Cummings, Ken Kolb, and David Sweet for helpful discussions. We are grateful for funding for this project from the Bradley University Sherry Endowment for Collaborative Student/Faculty Projects. The SEM and EDX studies were conducted at the University of Washington NanoTech User Facility (NTUF), a member of the NSF National Nanotechnology Infrastructure Network (NNIN). We would especially like to


_



thank Scott Braswell at the NTUF for assistance with these studies. Notes 1. That same day there were reports of car paint being damaged where it had come into contact with the particles. This fallout was not the first occurrence of car paint being damaged; however, it was never this severe in previous situations. At least 10 insurance claims were submitted for new car paint jobs (1). 2. Our efforts to demonstrate acid vapor attack on nylon in a closed system have produced erratic results. In these experiments, samples of nylon fabric (from nylon stockings) and nylon film (from oven cooking bags) were sprinkled with various oxides, sealed into plastic bags containing sulfur dioxide and water vapor, and sometimes exposed to various ultraviolet and visible light sources to simulate sunlight. Sometimes the nylon would degrade, sometimes it would not, producing tantalizing but inconsistent results. 3 Nylon swatches that had been rubbed on human sweat degraded, as did nylon that was soaked in deionized water for an hour or more, but nylon swatches that had been soaked in 5 g NaCl/100 mL water did not degrade.

Literature Cited 1. Kenyon, Theo J. Downtown Fallout Creates Slow Burn. Peoria Journal Star, June 11, 1970, p D16. 2. Sweet, D. Bradley University, Peoria, IL. Personal Interviews, 2006, 2007. 3. Jones, A. V.; Clemmet, M.; Higton, A.; Golding, E. Access to Chemistry; Royal Society of Chemistry: Cambridge, U.K., 1995; p 249. 4. Rayner-Canham, G.; Overton, T. Descriptive Inorganic Chemistry, 4th ed.; W. H. Freeman and Company: New York, 2006; pp 426-427. 5. Srivastava, R. K.; Miller, C. A.; Erickson, C.; Jambhekar, R. Emissions of Sulfur Trioxide From Coal-Fired Power Plants. Presented at POWER-GEN International 2000, Orlando, FL, 2002, http:// www.babcockpower.com/pdf/t-178.pdf (accessed Jan 2011). 6. Snyder, C. H. Chemicals, Pollution, and the Environment. The Extraordinary Chemistry of Ordinary Things, 2nd ed.; John Wiley and Sons, Inc.: New York, 1992; pp 352-353. 7. Nojima, K.; Yamaashi, Y. Chem. Pharm. Bull. 2004, 52, 335–338. 8. Finlayson-Pitts, B. J.; Pitts, J. N. Upper and Lower Atmosphere; Academic Press: San Diego, CA, 2000; pp 298-299, 325. 9. vanLoon, G. W.; Duffy, S. J. Environmental Chemistry: A Global Perspective; Oxford University Press: New York, 2000; p 100. 10. Scheetz, B. E. Chemistry and Mineralogy of Coal Fly Ash: Basis for Beneficial Use. Presented at Proceedings of State Regulation of Coal Combustion By-Product Placement at Mine Sites: A Technical Interactive Forum, Harrisburg, PA, 2004; pp 35-42. http://www. mcrcc.osmre.gov/MCR/Resources/ccb/PDF/State_Regulation_of_ CCB_Placement.pdf (accessed Jan 2011). 11. Craver, J. K.; Tess, R. W. Applied Polymer Science; Organic Coatings and Plastics Chemistry Division of The American Chemical Society: Washington, DC, 1975; pp 430-432.


_

12. Raber, L.Chem. Eng. NewsApril 12, 2010, p 46. 13. Keenan, C. W.; Kleinfelter, D. C.; Wood, J. S. General College Chemistry, 6th ed. Harper & Row: New York, 1980. 14. Zeronian, S. H.; Alger, K. W.; Omaye, S. T. Text. Res. J. 1973, 43, 228–237. 15. Brown, L.; Bul, V. T.; Bonin, H. W. ANTEC. 2004, 62, 3787– 3790. 16. Goss, L. M. J. Chem. Educ. 2003, 80, 39. 17. Epp, D. N.; Curtright, R. J. Chem. Educ. 1991, 68, 1034–1035. 18. Schilling, A. L.; Leber, P. A.; Yoder, C. H. J. Chem. Educ. 2009, 86, 225–226. 19. Schultz, R. F. Science Teaching Experiments: Chapter III. Environmental Experiments...from Edison: Edison Electric Institute: Washington, DC, available online at Charles Edison Fund, Science Teaching Experiments, http://www.charlesedisonfund.org/Experiments/HTMLexperiments/Chapter3/3-Expt2/p1.html (accessed Jan 2011). 20. Charola, A. E. J. Chem. Educ. 1987, 64, 436–437. 21. Baedecker, P. A.; Reddy, M. M. J. Chem. Educ. 1993, 68, 1034– 1035. 22. Shilling, A. L.; Hess, K. R.; Leber, P. A.; Yoder, C. H. J. Chem. Educ. 2004, 81, 274–277. 23. Mansfield, J. Utility Firms Have Problems Fighting Air Pollution. Peoria Journal Star, April 23, 1967; p A1. 24. Sax, N. I. Dangerous Properties of Industrial Materials, 4th ed.; Van Nostrand Reinhold Company: New York, 1975; p 1133. 25. Cronan, C. S.; Schofield, C. L. Science 1979, 204, 304–306. 26. Clean Air Act of 1970. Public Law 91-604, 1970. 27. Johnson, N. Ameren Corporation. Private communication, 2009. 28. Labinsky, D. Central Illinois Environment Better but There's Still Room To Improve. Peoria Journal Star, April 22, 1990, p A12. 29. Illinois Environmental Protection Agency. Illinois Annual Air Quality Report. http://www.epa.state.il.us/air/air-quality-report/ index.html (accessed Jan 2011). 30. Jones, A.; Slingo, A.; Ravishankara, A. R.; Liss, P. S.; Wolff, E.; Shine, K. P. Philis. Trans. R. Soc., B 1997, 352, 221–229. 31. Cox, R. A. Philis. Trans. R. Soc., B 1997, 352, 251–254. 32. Kunzig, R. Sci. Am. 2008, 299, 46–55. 33. Chemithon Enterprises, Environmental Systems: Flue Gas Conditioning Systems. http://www.chemithon.com/Enviro_fluegas. html (accessed Jan 2011). 34. Document NO. 5-04-0602 in the Appellate Court of Illinois Fifth District. http://www.state.il.us/court/opinions/appellatecourt/2006/ 5thdistrict/march/html/5040602.htm (accessed Jan 2011). 35. Bradsher, K.; Barboza, D.. Pollution from Chinese Coal Casts a Global Shadow. The New York Times, June 11, 2006. 36. Carmichael, G. R.; Streets, D. G.; Calori, G.; Amann, M.; Jacobson, M. Z.; Hansen, J.; Ueda, H. Environ. Sci. Technol. 2002, 36, 4707– 4713. 37. Zhang, L.; Jacob, D. J.; Kopacz, M.; Henze, D. K.; Singh, K.; Jaffe, D. A. Geophys. Res. Lett. 2009, 36, L11810.

Supporting Information Available A Microsoft Word document containing a scanning electron microscope (SEM) image and electron dispersion of X-rays (EDX) analysis of a recent sample of fly ash from a coal-burning power plant. This material is available via the Internet at http://pubs.acs.org.


_


_


391

Classroom Illustrations of Acidic Air Pollution Using ... - ACS Publications

Recommend Documents