- Acid Rain (mrnerix) Mac Os X
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'Acid rain' is a popular term referring to the deposition of a mixture from wet (rain, snow, sleet, fog, cloudwater, and dew) and dry (acidifying particles and gases) acidic components. Distilled water, once carbon dioxide is removed, has a neutral pH of 7. Liquids with a pH less than 7 are acidic, and those with a pH greater than 7 are alkaline. Make your day orange! Over 10 million iOS downloads! Join the party!:) - also available for iPad and Mac OS! A deadly genetically modified rain cloud of fruit dissolving vapour is heading for the orange capital of the world! A fruit loving, natural-ingredient pioneering he. LimeRa1n needs to be launched a certain way on Mac OS X and Windows. If you are running LimeRa1n on Windows you will need to ensure you grant it administrative privileges by right clicking it and selecting the ‘run as administrator option'. On Mac OS X 10.8 Mountain Lion, because of the new Gatekeeper software, there is also a trick to.
Inorganic Reactions Experiment
Authors: Rachel Casiday and Regina Frey |
Natural Acidity of Rainwater
Pure water has a pH of 7.0 (neutral); however, natural,unpolluted rainwater actually has a pH of about 5.6(acidic).[Recall from Experiment 1 that pH is a measure of thehydrogen ion (H+) concentration.] The acidity ofrainwater comes from the natural presence of three substances (CO2,NO, and SO2) found in the troposphere (the lowestlayer of the atmosphere). As is seen in Table I, carbon dioxide(CO2) is present in the greatest concentration andtherefore contributes the most to the natural acidity ofrainwater.
Gas | Natural Sources | Concentration |
| Carbon dioxide CO2 | Decomposition | 355 ppm |
| Nitric oxide NO | Electric discharge | 0.01 ppm |
| Sulfur dioxide SO2 | Volcanic gases | 0-0.01 ppm |
Table 1Carbon dioxide, produced in the decomposition of organic material, is the primary source of acidity in unpolluted rainwater. NOTE: Parts per million (ppm) is a common concentration measure used in environmental chemistry. The formula for ppm is given by: |
Carbon dioxide reacts with water to form carbonic acid(Equation 1). Carbonic acid then dissociates to give the hydrogenion (H+) and the hydrogen carbonate ion (HCO3-)(Equation 2). Social hunter mac os. The ability of H2CO3 todeliver H+ is what classifies this molecule as anacid, thus lowering the pH of a solution.
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Nitric oxide (NO), which also contributes to the naturalacidity of rainwater, is formed during lightning storms by thereaction of nitrogen and oxygen, two common atmospheric gases(Equation 3). In air, NO is oxidized to nitrogen dioxide (NO2)(Equation 4), which in turn reacts with water to give nitric acid(HNO3) (Equation 5). This acid dissociates in water toyield hydrogen ions and nitrate ions (NO3-)in a reaction analagous to the dissociation of carbonic acidshown in Equation 2, again lowering the pH of the solution.
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Acidity of Polluted Rainwater
Unfortunately, human industrial activity produces additionalacid-forming compounds in far greater quantities than the naturalsources of acidity described above. In some areas of the UnitedStates, the pH of rainwater can be 3.0 or lower, approximately1000 times more acidic than normal rainwater. Peel the potatoes mac os. In 1982, the pH ofa fog on the West Coast of the United States was measured at 1.8!When rainwater is too acidic, it can cause problems ranging fromkilling freshwater fish and damaging crops, to eroding buildingsand monuments.
Questions on Acidity of Rainwater
1. List two or more ways that you could test the acidity of asample of rainwater.
2. Write a balanced chemical equation for the dissociation ofnitric acid in water. (HINT: Draw an analogy with Equation 2.)
3. The gaseous oxides found in the atmosphere, including CO2and NO are nonmetal oxides. What would happen to the pH ofrainwater if the atmosphere contained metal oxides instead?(HINT: Think back to Experiment 1.) Briefly, explain your answer.
Sources of Excess Acidity in Rainwater
What causes such a dramatic increase in the acidity of rainrelative to pure water? The answer lies within the concentrationsof nitric oxide and sulfur dioxide in polluted air. As shown inTable II and Figure 1, the concentrations of these oxides aremuch higher than in clean air.
Gas | Non-Natural Sources | Concentration |
| Nitric oxide NO | Internal Combustion | 0.2 ppm |
| Sulfur dioxide SO2 | Fossil-fuel Combustion | 0.1 - 2.0 ppm |
Table IIHumans cause many combustion processes that dramatically increase the concentrations of acid-producing oxides in the atmosphere. Although CO2 is present in a much higher concentration than NO and SO2, CO2 does not form acid to the same extent as the other two gases. Thus, a large increase in the concentration of NO and SO2 significantly affects the pH of rainwater, even though both gases are present at much lower concentration than CO2. |
Figure 1Comparison of the concentrations of NO and SO2 in clean and polluted air. |
About one-fourth of the acidity of rain is accounted for bynitric acid (HNO3). In addition to the naturalprocesses that form small amounts of nitric acid in rainwater,high-temperature air combustion, such as occurs in car enginesand power plants, produces large amounts of NO gas. This gas thenforms nitric acid via Equations 4 and 5. Thus, a process thatoccurs naturally at levels tolerable by the environment can harmthe environment when human activity causes the process (e.g.,formation of nitric acid) to occur to a much greater extent.
What about the other 75% of the acidity of rain? Most isaccounted for by the presence of sulfuric acid (H2SO4)in rainwater. Although sulfuric acid may be produced naturally insmall quantities from biological decay and volcanic activity(Figure 1), it is produced almost entirely by human activity,especially the combustion of sulfur-containing fossil fuels inpower plants. When these fossil fuels are burned, the sulfurcontained in them reacts with oxygen from the air to form sulfurdioxide (SO2). Combustion of fossil fuels accounts forapproximately 80% of the total atmospheric SO2 in theUnited States. The effects of burning fossil fuels can bedramatic: in contrast to the unpolluted atmospheric SO2concentration of 0 to 0.01 ppm, polluted urban air can contain0.1 to 2 ppm SO2, or up to 200 times more SO2!Sulfur dioxide, like the oxides of carbon and nitrogen, reactswith water to form sulfuric acid (Equation 6).
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Sulfuric acid is a strong acid, so it readilydissociates in water, to give an H+ ion and an HSO4-ion (Equation 7). The HSO4- ion may furtherdissociate to give H+ and SO42-(Equation 8). Thus, the presence of H2SO4causes the concentration of H+ ions to increasedramatically, and so the pH of the rainwater drops to harmfullevels.
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Questions on Sources of Acidity in Rainwater
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Nitric oxide (NO), which also contributes to the naturalacidity of rainwater, is formed during lightning storms by thereaction of nitrogen and oxygen, two common atmospheric gases(Equation 3). In air, NO is oxidized to nitrogen dioxide (NO2)(Equation 4), which in turn reacts with water to give nitric acid(HNO3) (Equation 5). This acid dissociates in water toyield hydrogen ions and nitrate ions (NO3-)in a reaction analagous to the dissociation of carbonic acidshown in Equation 2, again lowering the pH of the solution.
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Acidity of Polluted Rainwater
Unfortunately, human industrial activity produces additionalacid-forming compounds in far greater quantities than the naturalsources of acidity described above. In some areas of the UnitedStates, the pH of rainwater can be 3.0 or lower, approximately1000 times more acidic than normal rainwater. Peel the potatoes mac os. In 1982, the pH ofa fog on the West Coast of the United States was measured at 1.8!When rainwater is too acidic, it can cause problems ranging fromkilling freshwater fish and damaging crops, to eroding buildingsand monuments.
Questions on Acidity of Rainwater
1. List two or more ways that you could test the acidity of asample of rainwater.
2. Write a balanced chemical equation for the dissociation ofnitric acid in water. (HINT: Draw an analogy with Equation 2.)
3. The gaseous oxides found in the atmosphere, including CO2and NO are nonmetal oxides. What would happen to the pH ofrainwater if the atmosphere contained metal oxides instead?(HINT: Think back to Experiment 1.) Briefly, explain your answer.
Sources of Excess Acidity in Rainwater
What causes such a dramatic increase in the acidity of rainrelative to pure water? The answer lies within the concentrationsof nitric oxide and sulfur dioxide in polluted air. As shown inTable II and Figure 1, the concentrations of these oxides aremuch higher than in clean air.
Gas | Non-Natural Sources | Concentration |
| Nitric oxide NO | Internal Combustion | 0.2 ppm |
| Sulfur dioxide SO2 | Fossil-fuel Combustion | 0.1 - 2.0 ppm |
Table IIHumans cause many combustion processes that dramatically increase the concentrations of acid-producing oxides in the atmosphere. Although CO2 is present in a much higher concentration than NO and SO2, CO2 does not form acid to the same extent as the other two gases. Thus, a large increase in the concentration of NO and SO2 significantly affects the pH of rainwater, even though both gases are present at much lower concentration than CO2. |
Figure 1Comparison of the concentrations of NO and SO2 in clean and polluted air. |
About one-fourth of the acidity of rain is accounted for bynitric acid (HNO3). In addition to the naturalprocesses that form small amounts of nitric acid in rainwater,high-temperature air combustion, such as occurs in car enginesand power plants, produces large amounts of NO gas. This gas thenforms nitric acid via Equations 4 and 5. Thus, a process thatoccurs naturally at levels tolerable by the environment can harmthe environment when human activity causes the process (e.g.,formation of nitric acid) to occur to a much greater extent.
What about the other 75% of the acidity of rain? Most isaccounted for by the presence of sulfuric acid (H2SO4)in rainwater. Although sulfuric acid may be produced naturally insmall quantities from biological decay and volcanic activity(Figure 1), it is produced almost entirely by human activity,especially the combustion of sulfur-containing fossil fuels inpower plants. When these fossil fuels are burned, the sulfurcontained in them reacts with oxygen from the air to form sulfurdioxide (SO2). Combustion of fossil fuels accounts forapproximately 80% of the total atmospheric SO2 in theUnited States. The effects of burning fossil fuels can bedramatic: in contrast to the unpolluted atmospheric SO2concentration of 0 to 0.01 ppm, polluted urban air can contain0.1 to 2 ppm SO2, or up to 200 times more SO2!Sulfur dioxide, like the oxides of carbon and nitrogen, reactswith water to form sulfuric acid (Equation 6).
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Sulfuric acid is a strong acid, so it readilydissociates in water, to give an H+ ion and an HSO4-ion (Equation 7). The HSO4- ion may furtherdissociate to give H+ and SO42-(Equation 8). Thus, the presence of H2SO4causes the concentration of H+ ions to increasedramatically, and so the pH of the rainwater drops to harmfullevels.
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Questions on Sources of Acidity in Rainwater
4. At sea level and 25oC,one mole of air fills a volume of 24.5 liters, and the density ofair is 1.22x10-6 g/ml. Compute the mole fraction (i.e.,moles of component /total moles) and molarity of SO2 when the atmosphericconcentration of SO2 is 2.0 ppm (see note in Table I).
5.One strategy for limiting theamount of acid pollution in the atmosphere is scrubbing.In particular, calcium oxide (CaO) is injected into thecombustion chamber of a power plant, where it reacts with thesulfur dioxide produced, to yield solid calcium sulfite.
a. Write a balanced chemical equation for this reaction. (HINT: Consult the table of common ions in the tutorial assignment for Experiment 1 to view the structure and formula for sulfite; also, use your knowledge of the periodic table to deduce the charge of the calcium ion. Using these facts, you can deduce the formula for calcium sulfite.)
Acid Rain (mrnerix) Mac Os X
b. Approximately one ton, or 9.0x102 kg, of calcium sulfite is generated each year for every person served by a power plant. How much sulfur dioxide (in moles) is prevented from entering the atmosphere when this much calcium sulfite is generated? Show your calculation.
c. The final stage in the scrubbing process is to treat the combustion gases with a slurry of solid CaO in water, in order to trap any remaining SO2 and convert it to calcium sulfite. A slurry is a thick suspension of an insoluble precipitate in water. Using the solubility guidelines provided in the lab manual for this experiment, predict whether this stage of the scrubbing process will produce a slurry (i.e., precipitate) or a solution (i.e., no precipitate) of calcium sulfite .
d. If MgO, rather than CaO, were used for scrubbing, would the product of the final stage be a slurry or a solution of magnesium sulfite? (Assume that a very large quantity of magnesium sulfite, relative to the amount of water, is produced.)
Acid Rain (mrnerix) Mac Os Download
Environmental Effects of Acid Rain
Acid rain triggers a number of inorganic and biochemicalreactions with deleterious environmental effects, making this agrowing environmental problem worldwide.
- Many lakes have become so acidic that fish cannot live in them anymore.
- Degradation of many soil minerals produces metal ions that are then washed away in the runoff, causing several effects:
- The release of toxic ions, such as Al3+, into the water supply.
- The loss of important minerals, such as Ca2+, from the soil, killing trees and damaging crops.
- Atmospheric pollutants are easily moved by wind currents, so acid-rain effects are felt far from where pollutants are generated.
Stone Buildings and Monuments in Acid Rain
Marble and limestone have long been preferred materials forconstructing durable buildings and monuments. The Saint Louis ArtMuseum, the Parthenon in Greece, the Chicago Field Museum, andthe United States Capitol building are all made of thesematerials. Marble and limestone both consist of calcium carbonate(CaCO3), and differ only in their crystallinestructure. Limestone consists of smaller crystals and is moreporous than marble; it is used more extensively in buildings.Marble, with its larger crystals and smaller pores, can attain ahigh polish and is thus preferred for monuments and statues.Although these are recognized as highly durable materials,buildings and outdoor monuments made of marble and limestone arenow being gradually eroded away by acid rain.
How does this happen? A chemical reaction (Equation 9) betweencalcium carbonate and sulfuric acid (the primary acid componentof acid rain) results in the dissolution of CaCO3 togive aqueous ions, which in turn are washed away in the waterflow.
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This process occurs at the surface of thebuildings or monuments; thus acid rain can easily destroy thedetails on relief work (e.g., the faces on a statue),but generally does not affect the structural integrity of thebuilding. The degree of damage is determined not only by theacidity of the rainwater, but also by the amount of water flowthat a region of the surface receives. Regions exposed to directdownpour of acid rain are highly susceptible to erosion, butregions that are more sheltered from water flow (such as undereaves and overhangs of limestone buildings) are much betterpreserved. The marble columns of the emperors Marcus Aurelius andTrajan, in Rome, provide a striking example: large volumes ofrainwater flow directly over certain parts of the columns, whichhave been badly eroded; other parts are protected by wind effectsfrom this flow, and are in extremely good condition even afternearly 2000 years!
Even those parts of marble and limestonestructures that are not themselves eroded can be damaged by thisprocess (Equation 9). When the water dries, it leaves behind theions that were dissolved in it. When a solution containingcalcium and sulfate ions dries, the ions crystallize as CaSO4l 2H2O,which is gypsum. Gypsum is soluble in water, so it is washed awayfrom areas that receive a heavy flow of rain. However, gypsumaccumulates in the same sheltered areas that are protected fromerosion, and attracts dust, carbon particles, dry-ash, and otherdark pollutants. This results in blackening of the surfaces wheregypsum accumulates.
An even more serious situationarises when water containing calcium and sulfate ions penetratesthe stone's pores. When the water dries, the ions form saltcrystals within the pore system. These crystals can disrupt thecrystalline arrangement of the atoms in the stone, causing thefundamental structure of the stone to be disturbed. If thecrystalline structure is disrupted sufficiently, the stone mayactually crack. Thus, porosity is an important factor indetermining a stone's durability.
Questions on Effects of Acid Rain
6. Based on the information described above aboutthe calcium ion, and the formula of calcium carbonate (CaCO3),deduce the charge of the carbonate ion. Also, in the structure ofthe carbonate ion, are any of the oxygens bonded to one another,or all the oxygens bonded to the carbon atom? (HINT: Consult thestructure of the common ions given in the tutorial for Experiment1).
7. In water, H2SO4 candissociate to yield two H+ ions and one SO42-ion. Write the net ionic equation for the reaction of calciumcarbonate and sulfuric acid. (See the introduction to Experiment2 in the lab manual for a discussion of net ionic equations.)
8. Which is a more durable building material,limestone or marble? Briefly, explain your reasoning.
Additional Links:
- The Environmental Protection Agency's site on acid rain presents the basics of this problem in an accessible format.
- The National Atmospheric Deposition Program features isopleth maps showing the concentrations of many different pollutants throughout the country.
- Another very interesting EPA site on acid rain explains the novel 'allowance trading system' strategy for getting companies to control their sulfur dioxide emissions.
References:
Acid Rain (mrnerix) Mac Os Downloads
Brown, Lemay, and Buster. Chemistry: the Central Science,7th ed. Upper Saddle River, NJ: Prentice Hall, 1997. p. 673-5.
Charola, A. 'Acid Rain Effects on Stone Monuments,' J.Chem. Ed.64 (1987), p. 436-7.
Petrucci and Harwood. General Chemistry: Principles andModern Applications, 7th ed. Upper Saddle River, NJ:Prentice Hall, 1997. p. 614-5.
Walk, M. F. and P.J. Godfrey. 'Effects of Acid Depositionon Surface Waters,' J. New England Water Works Assn.Dec. 1990, p. 248-251.
Zumdahl, S. Chem. Principles, 3rd ed. Boston:Houghton Mifflin, 1998. p. 174-6.
Stryer, L. Biochemistry, 4th ed., W.H. Freeman andCo., New York, 1995, p. 332-339.
Acknowledgements:
The authors thank Dewey Holten (Washington University) formany helpful suggestions in the writing of this tutorial.
The development of this tutorial was supported by a grant fromthe Howard Hughes Medical Institute, through the UndergraduateBiological Sciences Education program, Grant HHMI# 71192-502004to Washington University.
Copyright 1998, Washington University, All Rights Reserved.
Article Title
Authors
Abstract
Soils in the Adirondack Mountains are relatively young, having formed in sediments deposited from the last major glacial episode that ended in the Adirondacks about 12,000 years ago. Most of the region is covered by an acidic, sandy, low-fertility soil called a Spodosol. Over time, Spodosols develop distinct colorful horizons that can be clearly distinguished in soil pits and road cuts. Because Spodosols are naturally acidic and contain marginal concentrations of some elements necessary to sustain a healthy forested ecosystem, they are somewhat fragile and susceptible to chemical changes. A century's worth of acid deposition ('acid rain') has depleted these soils of some important nutrients and has mobilized aluminum, an element linked to fish mortality and forest decline. Acid rain has declined over the past several decades, and lakes, streams, and soils are beginning to recover. How long it will take is still unknown.
Recommended Citation
Acid Rain (mrnerix) Mac Os 11
April, Richard; Keller, Dianne; and Hluchy, Michele (2016) 'Soils and Soil Acidification in the Adirondack Mountains,' Adirondack Journal of Environmental Studies: Vol. 21 : No. 1 , Article 12.
Available at: https://digitalworks.union.edu/ajes/vol21/iss1/12
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