Nitrogen Oxides and Acid Rain

Introduction

Set up a small-scale demonstration of the production of nitrogen oxides that shows how quickly these gases interact with water and acidify it.

Concepts

• Acid rain
• Environmental chemistry
• pH
• Gases

Materials Included In Kit

Additional Materials Required

Background

Acid Rain and Nitrogen Oxides

The burning of fossil fuels produces two classes of gases that interact with water in the atmosphere to produce acid precipitation. Acid rain is the most well-known form of acid precipitation, but the acid pollutants can also cause damage in the forms of acid fog, acid snow, dew, and “dry deposition” (attached in the dry form to dust particles). The two classes of gases are sulfur oxides (SO_x) and nitrogen oxides (NO_x).

As conveyed by the name oxides, both of these gases are formed when oxygen chemically combines with the sulfur or the nitrogen. In the case of SO_x, the sulfur is present as a contaminant in the fuel. When the fuel burns, the sulfur combines with oxygen to produce SO_x. In the case of NO_x, it is the nitrogen in the air that combines with oxygen to form nitrogen oxides. Normally, atmospheric nitrogen and oxygen do not combine to form NOx. If they did, we would have a very acidic atmosphere since our air is 78% nitrogen and 21% oxygen. However, at high temperatures, nitrogen and oxygen react much more readily to form NO_x. When automobile engines burn gasoline or when electricity generating power plants burn fossil fuels (coal, oil and natural gas), the high temperatures of these combustion processes cause the nitrogen and oxygen in the air to combine to form NO_x.

Sulfur oxides are the better known of the two classes of gases involved in acid precipitation. Sulfur is present in varying degrees in coal and oil deposits. The general approach to reducing this aspect of the problem is to try to remove the sulfur from the fuel before combustion. In the northeastern United States, about two thirds of the acidic deposition comes from sulfuric acid formed from SO_x created by the burning of fossil fuels to produce electricity. Many of the power plants that are involved are located in the Midwest.

Since the prevailing wind currents are from West to East, the western U.S. does not have this problem of acidic pollution being transported there by wind. Instead, the major acidic pollutants are NO_x that are generated in the region by internal combustion vehicles—cars and trucks. Nitric acid, formed from NO_x, accounts for two-thirds of the acid deposition in California. It is much more difficult to reduce the formation of nitrogen oxides, since it is the air rather than the fuel which is the main source of the nitrogen. The NO_x pollution is increasing more rapidly than SO_x pollution.

As this demonstration dramatically indicates, these oxides readily interact with water to form acid. If your students are not familiar with acids and bases, it may help to explain that acidity has something to do with one of the fundamental properties of matter. The system we use for describing the relative acidity of different materials is the pH scale.

The pH scale is a logarithmic measuring system that is numbered from 0 to 14. A change in one unit actually represents a ten-fold change in amount. Going from pH 4 to pH 3 represents a 10-fold increase in acidity (the lower the number, the greater the acidity). A decrease in pH from 5 to 3 represents a hundred-fold increase in acidity. Absolutely pure water is neutral and has a pH value of 7. If a substance is on one side of neutral, (0 to 7) we have something that we call acidic. If it is on the other side of neutral, (7 to 14) we call it basic.

Rain and other forms of precipitation generally do not have a pH of 7 even in undisturbed natural conditions. Dissolved carbon dioxide, the same gas that gives soda water its effervescence and interesting taste, often lowers the pH to about 5.5. Life has adapted very well to this weakly acidic precipitation. In contrast, the SO_x and NO_x can make rainwater ten to a thousand times more acidic (as low as pH 2.5). Acid rain, or the more general term acid precipitation, refers to this much stronger acidification caused by the oxides of sulfur and nitrogen.

Acid precipitation causes damage in many ways. Most people have heard how acid rain can kill lakes. Because they live in water, aquatic organisms (ranging from microorganisms to fish) can be very sensitive to the pH of their liquid environment. Many lakes in sections of the United States, Canada and Europe have become so acidic that the organisms that used to flourish there have disappeared. But it isn’t just lakes that are damaged in this way. Soils can become too acidic or their essential minerals can be depleted as acid precipitation washes them away. The acids can also harm plants directly by damaging leaves and preventing seeds from germinating. Acid precipitation also dissolves and mobilizes trace metals in the soil such as aluminum and lead and can cause their concentration in the soil and in nearby lakes and streams to increase to toxic levels. On a more visible, though less life-threatening level, acids released into our atmosphere damage statues and erode artwork and buildings.

Demonstration

The nitrogen oxides are produced by the reaction of copper with nitric acid. The nitrogen dioxide that forms is bubbled into a large volume of water. By having a pH indicator present in the water reservoir, one can readily see that the gas acidifies the water.

The nitrogen dioxide that is produced is a brown gas. Since students see a brown gas forming and then bubbling into the water reservoir, it is important to use a pH indicator whose color change will not be confused with a possible coloring effect by the brown gas. Phenol red is a perfect indicator for this purpose. In the basic form it is red, and it turns yellow in the presence of acid. So students observe a brown gas bubbling into a red solution and turning it yellow.

This reaction setup has a dramatic and surprising finish. After all the bubbling stops, liquid from the reservoir is sucked back into the reaction flask and fills it practically to the top. The cause of this dramatic result is actually directly related to the phenomenon of acid rain. As the nitrogen oxides are formed, they drive out the air that was in the flask, and the flask becomes filled with NO_x. The reaction is exothermic (releases energy) so the flask becomes warm. After the reaction is complete, the flask begins to cool. As gas cools, it reduces in volume and this creates a partial vacuum that sucks some water from the reservoir back into the flask. The nitrogen dioxide gas in the flask readily dissolves in and reacts with this water. This causes a stronger vacuum in the flask since all of its gas has disappeared into the water. The result is that more water is sucked into the flask to replace the nitrogen dioxide gas, and the flask fills with water. Thus, this dramatic result shows again how readily nitrogen oxides interact with water.

The finish to the reaction is also dramatic because the liquid that fills the flask turns bluish. This is because the copper that is involved in the reaction has chemically changed from the solid metal form to the dissolved ionic salt form. Copper salts are generally blue in color. Students will also be surprised to see that the pennies have practically or even totally disappeared.

A good way to introduce the demonstration is to add phenol red (about 1 mL) to 200 mL of water in a flask. Add small amounts of vinegar and ammonia to show students the color changes from red (basic, ammonia) to orange (neutral) to yellow (acid, vinegar). Hold the flask against a white background (lab coats are great for this purpose). If the color is not deep enough, add more phenol red solution. Observing the sequence of color changes is a good way for students to begin to appreciate that the pH scale is a continuous spectrum ranging from strongly acidic, to moderate to weak acids, to neutral to weak bases, moderate and finally strong bases.

Wash out the flask. Add about 200 mL water and enough indicator to give a good color. If the water is yellow or orange, add just enough (one to a few) drops of dilute ammonia to turn it red. Then have a student continuously blow air through a straw into the flask until the water turns yellow. This demonstration introduces the concept that a gas can acidify water. In this case, the gas is carbon dioxide. Note that carbon dioxide is not a cause of acid rain. It only acidifies water a little bit (to pH 5.5). The SO_x and NO_x make water ten to a thousand times more acidic (as low as pH 2.5).

Safety Precautions

Nitric acid is severely corrosive; strong oxidant; toxic by inhalation; avoid contact with acetic acid and readily oxidized substances; avoid all body contact. Phenol red indicator solution is flammable. The nitrogen dioxide (NO₂) reddish brown gas is extremely toxic by inhalation; TLV 3 ppm in air; forms corrosive acids in contact with moisture/water; severely corrosive to eyes, skin and mucous membranes. Perform this demonstration under an efficiently operating fume hood. Wear chemical splash goggles, chemical-resistant apron and chemical-resistant gloves.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. Under the fume hood, carefully open the flasks to release any remaining gas. Dispose of the acidic solution according to Flinn Suggested Disposal Method #24b.

Procedure

Perform the following procedure under an efficiently operating fume hood.

1. Fill the large, narrow-mouth glass vessel about two-thirds full with water. Add enough phenol red solution (10 to 12 mL) to develop a reasonably deep color. If the color is orange or yellow, add just enough drops of dilute ammonia or sodium hydroxide to make the color red. Too much ammonia will cause soap bubbles.
2. Run the tubing from the side-arm flask into the large vessel so that the free end is at the bottom of the water in the large vessel. The gas bubbling from the side-arm flask needs to travel through as much water as possible so that it is absorbed rather than released into the air. Place Parafilm as securely as possible over the mouth of the large glass vessel without constricting the tubing.
3. Place two pre-1982 pennies or four to six grams of elemental copper in the bottom of the side arm flask.
4. Wearing chemical-resistant gloves, apron and face-shield or chemical splash goggles, carefully add 50 mL concentrated nitric acid to the side-arm flask. Immediately tightly stopper the side-arm flask.
5. Observe the gas bubbling into the large flask (first, the air from the side-arm flask comes through and then the brown nitrogen dioxide). Observe the change in the color of the water. Carefully agitate the side-arm flask to allow all the copper to react. Comment on the temperature at the bottom of the flask.
6. A minute or a few minutes after all the bubbling stops (be patient), observe the water rushing into the side-arm flask. Note the color of the water in the side-arm flask and the disappearance of the pennies or the copper.

Chemical Reactions

Cu(s) + 4HNO₃ → Cu(NO₃)₂ + 2NO₂ + 2H₂O

2NO₂ + H₂O → H⁺ + NO^3– + HNO₂

Student Worksheet PDF

10067_Student1.pdf

Further Extensions

Suggested Reading

Catalog of a variety of materials that are available from The Acid Rain Foundation, Inc., 1410 Varsity Drive, Raleigh, NC 27606.

Acid Rain—Teacher’s Guide from GEMS, Lawrence Hall of Science, University of California: Berkeley, CA 94720.

Acid Rain: The Chemistry of Acid Deposition from the Atmosphere from Institute for Chemical Education, Department of Chemistry, University of Wisconsin—Madison, 1101 University Avenue, Madison, WI 53706-1396.

Answers to Questions

1. Draw the setup of this demonstration. Then, describe what you observed.

Nitric acid was added to the side-arm flask, which contained two copper pennies. A red-brown gas is produced that moves through the rubber tubing to the flask containing water. The water inside turns from red to yellow, and a great deal of bubbling occurs. Once the bubbling had stopped for a few minutes, the water rushed into the side-arm flask and turned blue.

2. Write a balanced chemical equation for each reaction given below.

a. Copper pennies in nitric acid

Cu(s) + 4HNO₃ → Cu(NO₃)₂ + 2NO₂ + 2H₂O

b. Nitrogen dioxide and water

2NO₂ + H₂O → H⁺ + NO₃^– + HNO₂

3. Phenol red, an acid–base indicator, was present in the water flask. It is red in a base and yellow in an acid. Knowing this, describe the effect nitrogen dioxide had on the water.

When the nitrogen dioxide gas entered the water flask, the water quickly turned yellow. This shows that the nitrogen dioxide made the water highly acidic.

4. Name two harmful effects of acid rain. Then, name one way we can prevent acid rain in the first place.

Acid rain can lower the pH of lakes and other bodies of water so much that organisms can no longer survive in the environment. Acid rain can cause soils to be too acidic and deplete essential minerals which causes damage to the leaves of plants and inhibit seeds from germinating. We can help prevent acid rain by carpooling or using public transportation, thereby reducing the amount of gasoline that is produced and transmitted into the atmosphere.

Discussion

It is true for pH, as it is with temperature, that they are very broad scales, and life as we know it is comfortable only in a very narrow band or set of values. This is fairly obvious for the linear temperature scale that can range from absolute zero to many thousands of degrees. A swing of fifty degrees up or down marks the difference between the intolerable heat of the summer desert to the freezing cold of winter. A change in our body temperature of just a few degrees can be fatal.

Since the pH scale is logarithmic, this narrowness of the region that supports life is less obvious for acidity. We can approximate that many life forms can tolerate a range of pH between about 4.5 to 9.5. This seems to be about one-third of the entire pH range. However, this conclusion is misleading. Since each change in pH by one integer represents a ten-fold actual change, the range from pH 4.5 to 9.5 represents a range of l0⁵ variations in value. The entire pH scale, going from 0 to 14, represents a range of 10¹⁴ changes in amount. Therefore the five pH units that life can tolerate are only 10⁵ out of a possible 10¹⁴ values. Another way of saying this is that the part of the pH scale that life can tolerate is only 1 out of 10⁹, or one part per billion. Imagine a linear scale stretching from one end of the classroom to the other, with the most acid being at one end, neutral being in the middle and most basic at the other end. The part that life can tolerate would be less than one millionth of an inch wide.

An old American saying states, “An ounce of prevention is worth a pound of cure.” This saying almost always applies to environmental issues. The best way to solve a pollution problem is to avoid causing it rather than trying to reduce the impacts that occur after the pollutant has been produced. The best ways to avoid the pollution problems that are associated with our use of fossil fuels are to improve the efficiency of our energy use (conservation) and to utilize renewable energy sources that have fewer pollution impacts (solar, wind, hydro, biomass).

Many people have the mistaken impression that conservation means that we have to drastically reduce the comforts and conveniences that we currently enjoy. This is simply not true. If we drive a car that gets 50 miles per gallon and travel 1,000 miles, we will burn 20 gallons of gasoline and cause some pollution. If our car only gets 25 miles per gallon, the same trip will consume 40 gallons of gasoline and produce about twice the pollution. This is an example of a conservation or energy efficiency improvement that does not significantly change our lifestyle and yet can have a big effect on pollution. Other examples are lights and electrical appliances (refrigerators) that perform well and use less energy.

Some conservation measures do have noticeable impacts on our lives. For example, vast amounts of energy would be saved and pollution would decrease significantly if Americans carpooled and took public transportation rather than driving their single occupancy cars to work. Our society could improve its public transportation infrastructure so that it would become an attractive alternative to driving. This would be another example of a conservation practice that would reduce many pollutants, including the nitrogen oxides that contribute to acid rain.

References

This demonstration activity was developed and authored by Dr. Art Sussman, Far West Laboratory, San Francisco, CA 94107.