Tuesday, September 11, 2012

ACID RAIN AND GEOLOGY : (Teacher's guide)

Level: Upper elementary to senior high

Anticipated Learning Outcomes

Students will become aware of acid rain, an important environmental problem.
Students will learn about the pH scale.
Students will determine the acidity of their local rain water.
Students will learn a little about the types of rocks in their state/region.
Students will discover how the types of rocks in an area can mitigate the effects of acid rain.

Background: Given on student handout.


Clean, plastic containers (the number needed depends on how many rock types you choose--see item 6). Don't use glass containers. Acidic solutions can leach bases from glass; the bases will neutralize the acid solution.
Acidified water. Collect some rain water, which will naturally be acidified. If you can't collect rain water, you can add sulfuric acid to distilled water to pH 4 or 5. Diluted hydrochloric acid (HCl) or vinegar, which contains acetic acid (CH3COOH), could be substituted for sulfuric acid. Substitution of hydrochloric acid or vinegar will not change the observed results.
Some method to measure pH (short-range pH paper that covers pH 3 to 8, or a pH meter).
Stirring rods (popsicle sticks or plastic spoons should suffice).
Plastic wrap.
Three or four kinds of rocks, preferably crushed. Crushing exposes more surface area to the solution and thus speeds up the reaction. Try to choose rocks that are common in your state or region. One rock type, however, must be limestone, dolomite, or marble.
Geologic map of your state/region (optional). If the rock types you choose can be keyed to a geologic map, the exercise will be more relevant.

Procedure: Given on student handout.

Results and Discussion

Only the limestone, dolomite, or marble will react. This assumes that none of the other rock types selected contains carbonate minerals like calcite or dolomite. When placed in the acidic solution, the calcite or dolomite will effervesce; the reaction will be more apparent with calcite than with dolomite, more apparent with a strong acid than a weak one. The bubbles that are produced are carbon dioxide (a gas), one of the products of the reaction. As the calcite or dolomite continues to react, the pH of the solution should approach 7.

This kind of reaction may be familiar to those students who bake. Many recipes call for vinegar (acetic acid) and baking soda (sodium bicarbonate, a base). Students may recall that when the two are mixed, they bubble. Just as in the case of limestone and acid, the reaction between vinegar and baking soda produces carbon dioxide as one of the products. The neutralizing effect of baking soda makes this common household product an effective remedy for "acid indigestion."
All rock types that do not contain calcite or dolomite (or any of the less common carbonate minerals that you should not have to worry about) should not react with the acid, and thus should not effervesce or cause a change in the pH.
The main goal of this exercise, therefore, is to show that lakes or rivers in areas with calcite-rich rocks should not be in danger of acidification, even if the rain is very acidic. For example, dolomite-rich rocks occur in southern Wisconsin, but not in the northern part of the state; only the lakes in the north are in danger of being acidified. Have the students study any geologic maps you might have to determine other danger areas.
Another interesting consequence of acid rain can be brought into the discussion. Many buildings and statues are constructed of limestone or marble; concrete also typically contains calcite. Such buildings and statues in acid rain areas are slowly deteriorating because the gypsum (one of the products of the reaction) is 100 times more soluble than calcite, and therefore gets washed away with the rain water. In Europe and Asia, some structures have been standing for tens of centuries. Will they survive long under the threat of acid rain?



  1. The pH scale. Acidity of a solution is typically reported as a number that ranges from 0 to 14. The middle number, 7, designates a neutral solution, neither acidic nor basic. Numbers less than 7 refer to acidic solutions and higher numbers refer to basic solutions. The pH scale is not a linear scale; rather, it is logarithmic (base 10). For example, a change from 3 to 2 (a difference of just one) means that the resultant solution is ten times as acidic.

    The pH of some natural waters and common household solutions are depicted on the scale below. Note that unpolluted rain water is typically slightly acidic due to equilibration with atmospheric carbon dioxide (CO2).

  2. Acid rain. Acid rain is a complex problem that requires an interdisciplinary approach in order to understand its generation and potential harm to the environment. Ecologists, chemists, geologists, and climatologists all study various aspects of the problem.

    Natural processes of respiration, decay, lightning-induced forest fires, and volcanic eruptions release numerous compounds into the atmosphere. During the past century, human activities have been responsible for ever larger amounts of these compounds into the atmosphere, primarily through the burning of fossil fuels (oil, natural gas, and coal). We refer to these compounds as pollutants if their concentrations exceed or approach the tolerance limits of organisms. Once in the atmosphere, many of these compounds, whether natural or human made, go through complex chemical reactions that produce new compounds. Two of these new compounds are sulfuric acid (H2SO4) and nitric acid (HNO3). When it rains (or snows), these two acids are washed out of the atmosphere and, in the process, acidify the rain water. Hence, rain water falling from a polluted atmosphere is more acidic, typically between pH 4.1 and 5.6.

    Coal- and oil-burning industries, exhaust from automobiles, and some smelting industries emit many sulfur and nitrogen oxides that greatly contribute to acid rain. Because of the ever increasing quantity of sulfur and nitrogen oxides, rain falling from polluted atmospheres is becoming more and more acidic. In the United States, most industry can be found in the eastern part of the Midwest and in the Northeast. Therefore, one might expect acid rain problems in those areas. Not so simple. Climatic patterns modify the distribution of the pollutants, and thus the distribution of acid rain. For example, areas downwind from some industries might be in greater danger of acid rain than the areas immediately surrounding the industry that produced the air pollutants.

    When the acid rain falls, it is incorporated in our rivers, lakes, and soil. The acids in the rain cause numerous complex chemical reactions to occur, thus further modifying the river, lake, or soil chemistry. Many organisms can not live in acidified waters. For example, game fish such as trout, bass, and perch do not thrive in water below pH 5. Thus, acid rain causes a decrease in diversity of organisms, an ecologic problem. Because the numbers and types of game fish decrease, there is an economic consequence also.
  3. Effects of geology. If acid rain falls over a large area, why aren't all the lakes acidified? In part, the answer is geology. Some types of rocks can reduce (neutralize) the acidity of the rain, whereas other rocks have no effect. Calcite (CaCO3) and dolomite [CaMg(CO3)2] are two minerals that greatly mitigate the effects of acid rain; calcite and dolomite are the principal minerals that make up the rocks limestone and dolomite, respectively, as well as marble. For example, the case of sulfuric acid falling on limestone can be summarized by the following reaction:
H2SO4 + CaCO3 --> CaSO4.H2O + CO2

The sulfuric acid is neutralized as the mineral gypsum (CaSO4.H2O) and carbon dioxide are produced. Lakes located on, or rivers flowing through, limestone terrains will be neutralized; they will not suffer the consequences of acid rain. The same would be true in dolomite or marble terrains.

  1. Record selected rock types on the data sheet.
  2. Add about 125 ml (1/2 cup) of acidified solution to each plastic container.
  3. Record the pH of the starting solution above "time elapsed = 0" for each rock type on the data sheet.
  4. Slowly add about two teaspoonfuls of crushed rock to the acidified solution and stir with a stirring rod.
  5. Note any reactions or lack thereof.
  6. Loosely cover each container with plastic wrap to prevent evaporation of the water.
  7. At the end of the class period or at the end of the day (whichever best fits your school schedule), record the time elapsed and the pH.
  8. Record time elapsed and the pH on a daily basis until the pH no longer changes.

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