Classroom Illustrations of Acidic Air Pollution Using

r2011 American Chemical Society and Division of Chemical Education, Inc.

_pubs.acs.org/jchemeduc

_Vol. 88 No. 4 April 2011

_Journal of Chemical Education 387

10.1021/ed100604a Published on Web 02/14/2011

Chemistry for Everyone

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

June 10, 1970, seemed to be a fairly typical warm andmuggysummer day in central Illinois, but something out of the ordinaryarrived 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 ofthe coal-burning power plants. When lunchtime rolled around,the employees from the local stores walked to the local restau-rants to refuel for the rest of the afternoon. According to anarticle in the local paper, the “Peoria Journal Star” (1), the ladiesworking downtown noticed ash deposits and holes in their nylonstockings after returning from lunch. They concluded that theash had attacked the fibers of the nylon stockings. Many ladiesreturned their stockings to the store where they had beenpurchased, but the same disintegration occurred as they walkedback to work (1). When the employees returned to work, manyof them reported having headaches and burning of the eyes.1

The affected people were angered at the damage apparentlycaused by this ash, even though pollution regulations wereweaker in those days. Accusations were made against the powerplant, but a company representative denied the allegations bysaying, “We haven't changed anything we are doing that wouldcause this phenomenon, but we will investigate to see if there issomething that we aren't aware of that would have caused it” (1).The local news featured a company representative placing stock-ings under the smoke stack and no damage was observed (2). Thecompany claimed that because of this, there was no way that thesmoke could be blamed. Many people took this as fact. However,what happened next became part of the Bradley UniversityChemistry departmental lore. David Sweet, a student who wasresearching with Professor Tom Cummings in the ChemistryDepartment, believed that this demonstration was flawed. Theash that had fallen from the sky was fly ash, a byproduct of coalcombustion flying up and out the power plant smokestack. Itappeared to him that the fly ash falling from the smokestack hadcarried sulfur-containing oxyacids out of the exhaust plumedown to earth. When the ash came in contact with the nylonstockings, the acids attacked the fabric.

David Sweet was so infuriated by the power company'sclaims that he called both the company and a local TV station tocomplain. There was no success with the company, but a TVnews reporter asked him for an interview. Sweet went to the labto come up with a demonstration proving the company waswrong (2). He ran a sequence of experiments in which he passedsulfur dioxide from a small tank up a ceramic tube through nylonstockings in attempt tomimic the environmental conditions on theday the nylon damage occurred. The first experiment used a dry,room-temperature stocking, and as he passed the sulfur dioxidethrough it there was no damage. He then passed the sulfur dioxide

through a pair of stockings that were dampened by a slight aqueousmist from a spray bottle; still there was no damage. Sweet sprinkleda dry pair of stockings with transition-metal oxides, such as iron,chromium, and manganese oxides, and passed the sulfur dioxideand there was still no damage. Finally, Sweet dampened a pair ofstockings, sprinkled themwith transition-metal oxides, and passedthe sulfur dioxide through them, and they disintegrated. This reac-tion with the damp stockings showed how the weather and pollu-tants present in the air could be in the right balance to cause thedeterioration of the stockings (2). The TV news crew filmed thedemonstration and presented it to the public.

How the Events Happened

Industrial Production of Sulfuric Acid

Industrially, sulfur-containing oxyacids are produced on amassive scale. Sulfurous acid can be produced by first combiningsulfur and oxygen to produce sulfur dioxide (e.g., by burning sul-fur or heating sulfide ores in an excess of air):

SþO2 f SO2 ð1ÞThe sulfur dioxide can then be combined with water to

form sulfurous acid:H2Oþ SO2 f H2SO3 ð2Þ

Sulfuric acid is produced in a process called the contactprocess. Here, sulfur dioxide combines further with oxygen toform sulfur trioxide:

2SO2 þO2 h 2SO3 ð3ÞThe formation of sulfur trioxide is reversible and can be very

slow. For the industrial-scale production of sulfur trioxide, vana-dium(V) oxide is used to catalyze the oxidation reaction at about450 K (3, 4). Hypothetically, sulfuric acid could form by simplyreacting sulfur trioxide directly with water:

H2Oþ SO3 f H2SO4 ð4ÞHowever, this is a dangerously exothermic reaction. To

circumvent this, the sulfur trioxide is dissolved in concentratedsulfuric acid, which produces fuming sulfuric acid (also calledoleum). The fuming sulfuric acid can be more safely reacted withwater to make more quantities of concentrated sulfuric acid (4).

Sulfuric Acid from Coal Combustion

At coal-burning power plants, sulfur-containing com-pounds can be converted to sulfur-containing oxyacids by a num-ber of routes. There are also multiple potential sources of thissulfur. For example, iron sulfides such as marcasite and pyrite

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(both with the formula FeS2) can occur in coal and will burn toproduce iron oxides and sulfur oxides (5).

The conversion of sulfur dioxide to sulfur trioxide can occurwithin the flames of the boilers of the power plant, but thisexothermic reaction is not favored at these high temperatures (5).This reaction can also take place at somewhat lower temperatureswith the catalytic assistance of iron oxides on the fly ash and onthe surfaces of equipment in the plant (5, 6). Sulfur trioxide inthe flue gases combines with water vapor either in the flue gaseswithin the smokestack or with water vapor in the atmosphereoutside of the smokestack. At temperatures between about 370and 425 K, the sulfuric acid will condense in the air as droplets oron surfaces (5).

Sulfur dioxide can also be converted to sulfuric acid by otherroutes. Sulfur dioxide that leaves the smokestack of the powerplant can react with ultraviolet light in sunlight and other speciesin the air such as hydroxyl radicals, biacetyl, benzaldehyde, andnitrogen dioxide to produce sulfur trioxide, but whether thesulfur dioxide can be oxidized in sunlight without these speciespresent is controversial (7, 8). The sulfur dioxide can react withliquid water adsorbed onto the ash surfaces to produce sulfurousacid. The sulfite ions in this sulfurous acid solution can then beoxidized to sulfate ions to make sulfuric acid. This oxidationprocess can be catalyzed by iron(III) species in the solution; theash itself could act as a source of these ions (9). Hydrogen sulfitecan also be photochemically oxidized in an aqueous solutioncontaining solid iron(III) oxide (8).

Themetal oxides in the fly ash come frommetal compoundsin the coal that have reacted with the oxygen during the combus-tion process (10). Analysis of fly ash often yields combinations ofeight oxide components (6) in varying concentrations: SiO2,Al2O3, Fe2O3, CaO, MgO, Na2O, K2O, and SO3. The support-ing information contains a scanning electron microscope (SEM)image and electron dispersion of X-rays (EDX) analysis of arecent sample of fly ash from a coal-burning power plant.

Nylon

Nylon was first discovered and patented by Wallace Car-others and his research group at the DuPont Experimental Sta-tion (11, 12). Nylon is a polyamide, containing amide functionalgroups made by the condensation reactions of amine and

carboxylic acid functional groups. There are different types of nylonpolymers, but a common type used to make fibers for women'sstockings is nylon-6,6, which is made by polymerizing hexanedioicacid (C6H8O4) and 1,6-diaminohexane (C6H16N2) (13). The twomonomers each contain six carbon atoms, hence, the name of theproduct, nylon-6,6. This condensation reaction produces wateras the amide linkages are formed. The reverse reaction canoccur, resulting in the hydrolysis of the amide group by watermolecules in the presence of acid catalysts and heat to producethe monomers.

Studies have been performed to analyze the effect of acid onnylon. Research showed that nylon fibers degraded similarlywhen exposed to various concentrations of either hydrochloric orsulfuric acid at 50 �C (14). By placing the fibers in simulatedenvironmental conditions, researchers have shown that the mostsignificant damage to nylon fibers occurred when they were wetand exposed to light and 0.2 ppm sulfur dioxide. Other researchexplored the degradation of nylon fibers in varying acid condi-tions (ranging from distilled water to 1.0 M sulfuric acid) andtemperature conditions (20-90 �C). This research showed thatan increase in temperature increases the absorption of water bythe 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 andcorrosive sulfur dioxide, and we have had difficulties directlyreproducing the results from the sketchy details of this experi-ment. The nylon samples have not disintegrated in the short timescales that we have run the experiments, as we are unwilling tosend extensive quantities of sulfur dioxide up the flues of ourfume hoods. However, classroom or laboratory demonstrationscan be performed to illustrate portions of the overall chemicalevent. A sample of FeS2 (believed to be marcasite) found near acentral Illinois coal seam was ground and heated in a looselycorked 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 paperplaced into the tube turned red, indicating the presence of acidicvapor, most likely sulfur dioxide. Other methods of sulfurdioxide production involve the combination of sulfite or hydro-gen sulfite salts with acids to form sulfurous acid, which decom-poses to produce sulfur dioxide (16, 17). Perhaps the simplestway to produce sulfur dioxide is by burning sulfur in air (18).

A booklet that is available online describes simple environ-mental experiments, including one involving placing a nylon

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






stocking outdoors and inspecting it occasionally for holes thatmight be caused by acidic fly ash (19). To be an effective dem-onstration, the experiment seems to be limited to locations whereacidic ash would be present in sufficient concentrations to be anuisance. Other demonstrations of the effects of acid pollutionon the environment have been published. One study looked atthree major types of stone that are affected by acid rain: marble(limestone), sandstone, and granite, which are used frequently inmonuments and buildings (20). Laboratory runoff experimentshave been conducted to quantify the erosion of marble andlimestone by acid rain (21). Aquatic life is also affected by acidrain when it interacts with Al(OH)3 found in soil and clay,causing Al3þ runoff to be introduced into bodies of water whereit becomes harmful (22).

Demonstration 1

A more graphic (and quite simple) demonstration involvesdissolving 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 onthe stockings. The minimum concentration that seems to suc-cessfully dissolve the nylon stocking threads in a reasonable timeis 4 M. The higher the concentration of acid, the faster thethreads dissolve: 6 M acid works well, and concentrated sulfuricacid dissolves through the fabric, as illustrated in Figure 1. Itsometimes takes a little time for the fabric to dissolve. Therefore,it is recommended that when doing the demonstration, add acidto the fabric first and then explain the connection to the nylondegradation event in Peoria while waiting for dissolution to takeplace. The nylon fabric can change color in the vicinity of theholes dissolved in the fabric. The very low pH of the acid dropletsappears to shift the colors of the fabric dyes, much like acid-baseindicators. However, the specific color the fabric turns can be un-predictable; one tan-colored stocking has turned red, and an-other brand has turned blue, presumably because different dyeswere used to achieve the specific tan colors. This demonstrationcan be shown at a variety of grade levels and can be shown to largegroups with an overhead projector. If the demonstration is notperformed on an overhead projector, it is easier to see when thenylon fabric is placed over a color-contrasting background. Watercan be used to clean up the demonstration, but some surfaces canrequire a bit of scrubbing to remove the sticky, gummy, partially-dissolved nylon.

Demonstration 2

We have also developed a variation on Sweet's originalexperiment that does not require a sulfur dioxide gas tank. Thisopen system2 still produces some sulfur dioxide and must beperformed in a fume hood. To perform this experiment, shownin Figure 2, combine 7.0 g of sodium bisulfite and 200 mL ofwater in a 500 mL Florence flask. (We used ACS reagent gradesodium bisulfite, a mixture of NaHSO3 andNa2S2O5.)Wipe themouth of the flask dry to remove any possible chemical contam-ination. Stretch about four stacked ∼3 cm squares of dry (orwater-soaked3) nylon stocking fabric tightly over the mouth of theflask and then secure the fabric to the flask mouth with coated oruncoated wire. Sprinkle approximately 0.03 g of iron(III) oxidepowder onto the fabric. Some powder will likely fall through thefabric layers into the sodium bisulfite solution, but this does notaffect the reaction. Place the flask assembly on a hot plate to boil

the solution, releasing some warmth, water vapor, and sulfurdioxide up into the fabric at the mouth of the flask. Before thesolution in the flask boils, use a clamp to hold a 400 mL beakerupside down over nylon at the mouth of the flask to help keep anyvapors that pass through the nylon fabric layers in the vicinity ofthose layers. Tilt the beaker at an angle of roughly 20� from verti-cally upside down so that any condensation droplets that collectwithin the beakerwillmove along the interior beaker walls and notdrip onto the nylon fabric. Rather, the droplets will move downthe 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 solu-tion 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 mayappear later if the entire experimental setup is left to coolovernight. The reason for this might be due to be a slow reactionbetween the acid and the nylon fibers or it might be due to achange in acid concentration on the fibers as the moisture on thefabric dries out overnight. The threads appear to break most nearwhere the fabric meets the lip of the flask, where the threadscurve the most and are under the most stress. Different samplesof nylon fabric appear to have varying susceptibility to attack bythe vapors, but a lack of sodium bisulfite in solution (and there-fore no sulfur dioxide production) results in no nylon damage.Adding powdered iron(III) oxide to the nylon is much moredamaging than adding no oxide at all. It is hypothesized that theiron(III) oxide catalyzes the formation of sulfur trioxide (andtherefore sulfuric acid) at the nylon, increasing the fabric damage.Adding the iron(III) oxide directly to theNaHSO3 solution ratherthan the fabric did not produce degradation of the nylon. Iron(III)oxide has been the best catalyst for these experiments. Addingpowdered vanadium(V) oxide, used in industrial production of

Figure 2. Demonstration of nylon fabric degradation. The fabric issprinkled with iron(III) oxide powder and the solution in the Florenceflask contains sodium bisulfite.

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sulfur dioxide, does not typically produce nearly as significantdamage as does powdered iron(III) oxide and is also not recom-mended for classroom use due its toxicity. Fly ash from a localpower plant also does not typically produce nearly as significantdamage as does powdered iron(III) oxide. Powdered iron(III)oxide that has been heated and then cooled does not reproduciblyproduce more significant damage than just the powdered oxidefrom the reagent bottle and often produces less damage. A sampleof powdered FeS2 (believed to be marcasite) does appear to be ableto produce nylon damage. This powder can also help damagenylon after it has been preheated on a hot plate and then cooled,presumably producing iron oxides, before being placed on thefabric.

Hazards

Safety precautions such as eye and skin protection must beobserved. Sulfuric acid is very corrosive. Sodium bisulfite canproduce toxic or corrosive gases upon exposure to heat or acids.Iron(II) sulfide (marcasite or pyrite) can produce toxic orcorrosive gases upon exposure to heat or acids. Iron(III) oxideis fairly inert. However, the bright red powder can dust surfaces(clothing, countertops, sinks, etc.) quite effectively and can bedifficult to remove.

Discussion

Ultimately, there were no significant repercussions orreparations resulting from David Sweet's demonstration, eventhough the power company was shown to be in error. The TVstation footage of the demonstration has been lost to time. Thereason this mysterious stocking-damaging event is rather uniquein Peoria history is not well understood. Perhaps there was anextraordinary quantity of iron oxide and sulfur oxide produced inthe plant emissions that day. However, a 1967 article in the“Peoria Journal Star” (23) shows the aforementioned powercompany was aware of emissions problems at the downtownPeoria station, which was built in 1890.

The historical anecdote and accompanying demonstrationsmake dramatic illustrations of concepts of the acidity of non-metal oxides, catalytic behavior, and air pollution. Air pollutionwill continue as nations develop industrially but what is done toreduce the pollution will influence the health, wealth, and wellbeing of those nations. The effects of sulfur oxide pollution frompower plants extend beyond a few torn nylon stockings. It affectsthe air we breathe, leading to health problems ranging from skin,eye, and upper respiratory irritation in low concentrations (recallthe headaches and eye irritation experienced by passers-by thatday), to asthma, edema of the lungs, and even respiratory paralysisin higher concentrations (24). This is why sulfur oxide-producingreactions should be handled in a fume hood. Acid precipitationincreases the hydrogen ion and aluminum ion concentrations inwaterways, harming life forms in the water (25).

The Clean Air Act and its subsequent amendments (includ-ing one in 1970) set standards on the emission of air pollutionfrom urban, industrial, and motor vehicle sources to protect theair quality and public health of the United States (26). Powercompanies have responded in a variety of ways to these standards.The age of the downtown Peoria plant and the cost of rebuildingto acceptable standards led the power company to close thestation on May 2, 1971 (23, 27). For other power plants, theseresponses included installing scrubber systems and burning coal

with lower sulfur content. Though Peoria had many local coalmines in the past, some of the central Illinois coal-fired powerplants prefer to burn significant quantities of lower-sulfur coaldelivered by rail from the Powder River Basin in the stateof Wyoming (27). One of the open-pit coal mines in the basinis actually crossed by Interstate 90 west of Gillette, WY, over800 mi (1300 km) from Peoria. As a result of the sulfur oxideregulations, the quantity of sulfur dioxide emission in Peoriadecreased from 32 ppb in 1972 to 7 ppb in 1989 (28). The av-erage level of sulfur dioxide in Peoria county declined from 7 ppbin 1995 to 2 ppb in 2007 (29).

However, complicated issues associated with sulfur oxideemissions persist. For example, oxidation of sulfur dioxide to sul-fates can contribute to the growth of tiny water droplets in theatmosphere. These sulfate-containing aerosols can reflect sun-light away from the earth, producing a cooling effect (30-32).Some people have proposed deliberately adding sulfur dioxide tothe atmosphere to increase the degree of sulfate aerosol cooling ofthe earth in an effort to combat global warming (32).

Another example of the complexities of sulfur oxide emis-sion involves fly ash, which can be removed from power plantemissions by electrostatic precipitation. This process can be faci-litated by a sulfur trioxide flue gas conditioning (FGC) systemwhere sulfur dioxide is converted to sulfur trioxide and deliber-ately added to the flue gases in the chimney to make them moreelectrically conductive in order to capture more fly ash (8, 33).Selective catalytic reduction (SCR) systems, designed to decreasenitrogen oxide emissions from power plants, have sometimesassisted the conversion of sulfur dioxide to sulfur trioxide (5, 34).In 2004, a catalyst problem of this type caused a power plant inIndiana to produce sulfur oxide emissions that blew into thecommunity of Mt. Carmel, IL. Residents there encounteredphysical problems (e.g., eye irritation) similar to that encoun-tered 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 fromsulfur oxides is becoming more apparent as more nations such asIndia and China become increasingly industrialized. China cur-rently burns more coal for energy than the United States andEuropean Union combined and builds more coal-fired plants atthe rate of about one per week (34) tomeet the energy needs of itspopulation of over one billion. There have been complaints fromJapan and South Korea about increases in the concentration ofsulfur dioxide in their air as a result of cross-border contamina-tion from China (35, 36). United Sates satellites and ground-based detectors in California, Oregon, andWashington have alsodetected Asian pollutants wafting into North America fromacross the Pacific Ocean (35, 37). Sulfur oxide acidification ofthe environment has been and will continue to be an issue forsome time.

Acknowledgment

We would like to thank Robert Gayhart, Max Taylor,Thomas Cummings, Ken Kolb, and David Sweet for helpfuldiscussions. We are grateful for funding for this project fromthe Bradley University Sherry Endowment for CollaborativeStudent/Faculty Projects. The SEM and EDX studies were con-ducted at the University ofWashingtonNanoTechUser Facility(NTUF), a member of the NSF National NanotechnologyInfrastructure Network (NNIN). We would especially like to






thank Scott Braswell at the NTUF for assistance with thesestudies.

Notes

1. That same day there were reports of car paint being damagedwhere it had come into contact with the particles. This falloutwas not the first occurrence of car paint being damaged;however, it was never this severe in previous situations. At least10 insurance claims were submitted for new car paint jobs (1).

2. Our efforts to demonstrate acid vapor attack on nylon in aclosed system have produced erratic results. In these experi-ments, samples of nylon fabric (from nylon stockings) andnylon film (from oven cooking bags) were sprinkled withvarious oxides, sealed into plastic bags containing sulfur dioxideand water vapor, and sometimes exposed to various ultravioletand visible light sources to simulate sunlight. Sometimes thenylon would degrade, sometimes it would not, producingtantalizing but inconsistent results.

3 Nylon swatches that had been rubbed on human sweat de-graded, as did nylon that was soaked in deionized water for anhour or more, but nylon swatches that had been soaked in 5 gNaCl/100 mL water did not degrade.

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Supporting Information Available

A Microsoft Word document containing a scanning electronmicroscope (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.

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Classroom Illustrations of Acidic Air Pollution Using