General Chemistry Labs

Acid Rain
Inorganic Reactions Experiment

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Authors: Rachel Casiday and Regina Frey
Department of Chemistry, Washington University
St. Louis, MO 63130

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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.


Natural Sources


Carbon dioxide
355 ppm
Nitric oxide
Electric discharge
Lightning Discharge
0.01 ppm
Sulfur dioxide
Volcanic gases
Voclanic 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:

Acid Rain

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.

CO2 + H2O → H2CO3
H2CO3 → H+ + HCO3-

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-) again lowering the pH of the solution.

N2 + O2
NO(g) + ½O2(g) → NO2(g)
3NO2(g) + H2O → 2HNO3(aq) + NO(g)

pH Paper

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.


Non-Natural Sources


Nitric oxide
Internal Combustion Internal Combustion 0.2 ppm
Sulfur dioxide
Fossil-fuel Combustion 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.

Comparison of the concentrations of NO and SO2 in clean and polluted air.

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 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.

H2SO4 → HSO4- + H+
HSO4- → SO42- + H+

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.

  • 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 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.

CaCO3(s) + H2SO4(aq) → Ca2+(aq) + SO42- + H2O + CO2

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.

Additional Links:


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.


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.

Revised: 1/7/11