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Acid Rain
Inorganic Reactions Experiment
Authors: Rachel Casiday and Regina Frey
Department of Chemistry, Washington
University
St. Louis, MO 63130
For information or comments on this tutorial, please contact R. Frey at [email protected].
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 the
hydrogen ion (H+) concentration.] The acidity of rainwater comes from the natural
presence of three substances (CO2, NO, and SO2) found in the troposphere (the lowest
layer of the atmosphere). As is seen in Table I, carbon dioxide (CO2) is present in the
greatest concentration and therefore contributes the most to the natural acidity of
rainwater.
Gas
Natural Sources
Carbon dioxide
Decomposition
CO2
Concentration
355 ppm
Electric
Nitric oxide
NO
discharge
Sulfur dioxide
SO2
Volcanic gases
0.01 ppm
0-0.01 ppm
Table 1
Carbon 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 hydrogen ion (H+) and the hydrogen carbonate ion (HCO3-)
(Equation 2). The ability of H2CO3 to deliver H+is what classifies this molecule as an
acid, thus lowering the pH of a solution.
(1)
(2)
Nitric oxide (NO), which also contributes to the natural acidity of rainwater, is formed
during lightning storms by the reaction of nitrogen and oxygen, two common
atmospheric gases (Equation 3). In air, NO is oxidized to nitrogen dioxide (NO 2)
(Equation 4), which in turn reacts with water to give nitric acid (HNO3) (Equation 5).
This acid dissociates in water to yield hydrogen ions and nitrate ions (NO 3-) in a reaction
analagous to the dissociation of carbonic acid shown in Equation 2, again lowering the
pH of the solution.
(3)
(4)
(5)
Acidity of Polluted Rainwater
Unfortunately, human industrial activity produces additional acid-forming compounds in
far greater quantities than the natural sources of acidity described above. In some areas
of the United States, the pH of rainwater can be 3.0 or lower, approximately 1000 times
more acidic than normal rainwater. In 1982, the pH of a fog on the West Coast of the
United States was measured at 1.8! When rainwater is too acidic, it can cause problems
ranging from killing freshwater fish and damaging crops, to eroding buildings and
monuments.
Questions on Acidity of Rainwater
1. List two or more ways that you could test the acidity of a sample of rainwater.
2. Write a balanced chemical equation for the dissociation of nitric acid in water. (HINT:
Draw an analogy with Equation 2.)
3. The gaseous oxides found in the atmosphere, including CO 2 and NO are nonmetal
oxides. What would happen to the pH of rainwater 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 rain relative to pure water? The
answer lies within the concentrations of nitric oxide and sulfur dioxide in polluted air.
As shown in Table II and Figure 1, the concentrations of these oxides are much higher
than in clean air.
Gas
Non-Natural
Sources
Nitric oxide
NO
Internal Combustion
Sulfur dioxide
SO2
Fossil-fuel
Concentration
0.2 ppm
0.1 - 2.0 ppm
Combustion
Table II
Humans 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 1
Comparison of the concentrations of
NO and SO2 in clean and polluted air.
About one-fourth of the acidity of rain is accounted for by nitric acid (HNO3). In
addition to the natural processes that form small amounts of nitric acid in rainwater,
high-temperature air combustion, such as occurs in car engines and power plants,
produces large amounts of NO gas. This gas then forms nitric acid via Equations 4 and
5. Thus, a process that occurs naturally at levels tolerable by the environment can harm
the 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 is accounted for by the presence
of sulfuric acid (H2SO4) in rainwater. Although sulfuric acid may be produced naturally
in small 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 in power plants. When these fossil fuels are burned, the sulfur contained in them
reacts with oxygen from the air to form sulfur dioxide (SO2). Combustion of fossil fuels
accounts for approximately 80% of the total atmospheric SO2 in the United States. The
effects of burning fossil fuels can be dramatic: in contrast to the unpolluted atmospheric
SO2 concentration of 0 to 0.01 ppm, polluted urban air can contain 0.1 to 2 ppm SO 2, or
up to 200 times more SO2! Sulfur dioxide, like the oxides of carbon and nitrogen, reacts
with water to form sulfuric acid (Equation 6).
(6)
Sulfuric acid is a strong acid, so it readily dissociates in water, to give an H+ ion and an
HSO4- ion (Equation 7). The HSO4- ion may further dissociate to give H+ and SO42(Equation 8). Thus, the presence of H2SO4causes the concentration of H+ ions to
increase dramatically, and so the pH of the rainwater drops to harmful levels.
(7)
(8)
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 of
air is 1.22x10-6 g/ml. Compute the mole fraction (i.e., moles of component /total moles)
and molarity of SO2 when the atmospheric concentration of SO2 is 2.0 ppm (see note in
Table I).
5. One strategy for limiting the amount of acid pollution in the atmosphere
is scrubbing. In particular, calcium oxide (CaO) is injected into the combustion chamber
of a power plant, where it reacts with the sulfur 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.)
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.)
Environmental Effects of Acid Rain
Acid rain triggers a number of inorganic and biochemical reactions with deleterious
environmental effects, making this a growing environmental problem worldwide.

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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:
3+
o The release of toxic ions, such as Al , into the water supply.
2+
o The loss of important minerals, such as Ca , 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 for constructing durable
buildings and monuments. The Saint Louis Art Museum, the Parthenon in Greece, the
Chicago Field Museum, and the United States Capitol building are all made of these
materials. Marble and limestone both consist of calcium carbonate (CaCO3), and differ
only in their crystalline structure. Limestone consists of smaller crystals and is more
porous than marble; it is used more extensively in buildings. Marble, with its larger
crystals and smaller pores, can attain a high 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 are now being gradually eroded away
by acid rain.
How does this happen? A chemical reaction (Equation 9) between calcium carbonate
and sulfuric acid (the primary acid component of acid rain) results in the dissolution of
CaCO3 to give aqueous ions, which in turn are washed away in the water flow.
(9)
This process occurs at the surface of the buildings or monuments; thus acid rain can
easily destroy the details on relief work (e.g., the faces on a statue), but generally does
not affect the structural integrity of the building. The degree of damage is determined not
only by the acidity of the rainwater, but also by the amount of water flow that a region of
the surface receives. Regions exposed to direct downpour of acid rain are highly
susceptible to erosion, but regions that are more sheltered from water flow (such as
under eaves and overhangs of limestone buildings) are much better preserved. The
marble columns of the emperors Marcus Aurelius and Trajan, in Rome, provide a
striking example: large volumes of rainwater flow directly over certain parts of the
columns, which have been badly eroded; other parts are protected by wind effects from
this flow, and are in extremely good condition even after nearly 2000 years!
Even those parts of marble and limestone structures that are not themselves eroded can
be damaged by this process (Equation 9). When the water dries, it leaves behind the ions
that were dissolved in it. When a solution containing calcium and sulfate ions dries, the
ions crystallize as CaSO4 2H2O, which is gypsum. Gypsum is soluble in water, so it is washed
away from areas that receive a heavy flow of rain. However, gypsum accumulates in the same
sheltered areas that are protected from erosion, and attracts dust, carbon particles, dry-ash, and other
dark pollutants. This results in blackening of the surfaces where gypsum accumulates.
An even more serious situation arises when water containing calcium and sulfate ions penetrates the
stone's pores. When the water dries, the ions form salt crystals within the pore system. These
crystals can disrupt the crystalline arrangement of the atoms in the stone, causing the fundamental
structure of the stone to be disturbed. If the crystalline structure is disrupted sufficiently, the stone
may actually crack. Thus, porosity is an important factor in determining a stone's durability.
Questions on Effects of Acid Rain
6. Based on the information described above about the calcium ion, and the formula of
calcium carbonate (CaCO3), deduce the charge of the carbonate ion. Also, in the
structure of the carbonate ion, are any of the oxygens bonded to one another, or all the
oxygens bonded to the carbon atom? (HINT: Consult the structure of the common ions
given in the tutorial for Experiment 1).
7. In water, H2SO4 can dissociate to yield two H+ ions and one SO42- ion. Write the net
ionic equation for the reaction of calcium carbonate and sulfuric acid. (See the
introduction to Experiment 2 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:
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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:
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 and Modern Applications, 7th ed.
Upper Saddle River, NJ: Prentice Hall, 1997. p. 614-5.
Walk, M. F. and P.J. Godfrey. "Effects of Acid Deposition on 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 and Co., New York, 1995, p. 332-339.
Acknowledgements:
The authors thank Dewey Holten (Washington University) for many helpful suggestions
in the writing of this tutorial.
The development of this tutorial was supported by a grant from the Howard Hughes
Medical Institute, through the Undergraduate Biological Sciences Education program,
Grant HHMI# 71192-502004 to Washington University.
Copyright 1998, Washington University, All Rights Reserved.