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
|
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 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:
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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.
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 (NO2)
(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 (NO3-)
in a reaction analagous to the dissociation of carbonic acid
shown in Equation 2, again lowering the pH of the solution.
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.
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
|
Concentration
|
Nitric oxide
NO |
Internal Combustion |
0.2 ppm |
Sulfur dioxide
SO2 |
Fossil-fuel Combustion |
0.1 - 2.0 ppm |
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.
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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 SO2, 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).
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 H2SO4
causes the concentration of H+ ions to increase
dramatically, and so the pH of the rainwater drops to harmful
levels.
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.
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.
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 CaSO4l 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.