In the Classroom
Overhead Projector Demonstrations
Generating Water-Soluble Noxious Gases: An Overhead Projector Demonstration Sally Solomon,* Maria Oliver-Hoyo, and Chinhyu Hur Department of Chemistry, Drexel University, Philadelphia, PA 19104
The substitution of an alkane by a halogen has been demonstrated on the overhead projector using a special type of divided dish (called a Conway dish) to collect the hydrogen halide gas product (1). This round transparent dish consists of an inner compartment where the reaction is performed and an outer area containing water. The two compartments are separated by a wall whose height is lower than the outside wall of the dish. When the dish is covered, vapors can migrate from the inner reaction area to the water-filled outer area. Instead of using this specialized dish, a simple inexpensive apparatus for generating and collecting water-soluble gases can be made from two beakers and a Petri dish. One small beaker (30 or 50 mL) serves as the reaction vessel, which is placed on a 10-cm Petri dish. Water is poured into the Petri dish and as soon as the reactants that produce the gas are mixed, the larger beaker is immediately placed over the smaller beaker so that the gas cannot escape into the atmosphere. Some of the gas diffuses into the water, where it can be detected by adding an appropriate reagent. One advantage of overhead projector demonstrations such as these is that very small amounts of reagents are needed; hence disposal is convenient. Reactions that generate sulfur dioxide and nitrogen dioxide using this apparatus along with detection methods are described below. Variations using household products are suggested where possible. Sulfur Dioxide Sulfur dioxide gas is a toxic colorless gas with a choking odor. Sulfur dioxide is an irritant with a “hazard to health” rating of 4 for eye contact, inhalation, and skin irritation, on a scale of 1 (least hazardous) to 5 (most hazardous). The threshold limit value is 5 ppm or 13 mg/m3 of air (2). Since the solubility in water of SO2 is 5–50 g/100 mL (2), the amount generated in the reactions described below dissolves readily in the 30 mL of water contained in the outer Petri dish.
Detecting Sulfur Dioxide The SO2 is detected in several different ways. The simplest takes advantage of the acid-forming properties of sulfur dioxide, which, as it dissolves in the water in the Petri dish, causes the solution to become more and more acidic: SO2(aq)
HSO 3᎑(aq) + H+(aq)
The increasing concentration of gas in the water causes an indicator to turn color starting from the lip of the overturned *Corresponding author.
beaker and spreading throughout the water in the Petri dish. When Universal Indicator (Fisher Scientific, catalog no. SOI-60) is chosen the change from green (neutral) to yellow takes 2 or 3 minutes or less, progressing to a reddish color (acidic) after 3–5 minutes. A second Petri dish containing water and indicator may be placed on the overhead so that the viewer can recall the original color. Another detection method depends upon the ability of very small quantities of SO2 to form colorless complexes with anthocyanin dyes (3, 4). A deep purple concentrated anthocyanin extract is made by mixing one-half head of chopped red cabbage with 1 L of rubbing alcohol in a 2-L beaker, boiling for one hour, and reducing to about 50 mL. Ten drops of the concentrated red cabbage extract is added to the water in the Petri dish. As SO2 is generated the purple anthocyanin first turns pink as the pH drops, then gradually decolorizes owing to the formation of a colorless complex. The reaction below is shown for an anthocyanin aglycone core, the chemical structure without the attached carbohydrate (5). O
7
HO
OH
O
+ SO2
3 5
OH
OH
+
HO
OH
OH OH H SO3H
OH
OH
Cyanidin
colorless complex
Attachment of sugars to the hydroxyl groups in the 3, 5, and 7 positions provides variations of the structure found in different plants (6 ). A third way to detect the sulfur dioxide is by its reduction of deep violet permanganate, MnO4᎑, to form the pale pink, nearly colorless MnII. 2MnO4᎑ + H+ + 5SO2 + 2H2O → 2MnII + 5HSO 4᎑ The color change takes place within minutes.
Generating Sulfur Dioxide Sulfur dioxide can be generated as an accompaniment to discussions on the production and properties of sulfur dioxide. Depending upon the detection method chosen, topics such as acids and bases, the formation of complexes, or oxidation– reduction may also be demonstrated. Topics that correspond to the specific generating reaction are mentioned below. Sulfite or bisulfite salts added to a strong acid produce sulfur dioxide gas: SO32᎑(aq) + 2H+(aq)
SO2(g) + H2O
HSO3᎑(aq)
SO2(g) + H2O
+
+ H (aq)
JChemEd.chem.wisc.edu • Vol. 75 No. 12 December 1998 • Journal of Chemical Education
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In the Classroom
The following chemicals and equipment are needed: Na2SO 3 or NaHSO3 6 M H2SO4; CAUTION: be sure to check that the acid is sufficiently concentrated Detection reagent (universal indicator, 0.002 M KMnO4, or anthocyanin) Overhead assembly (10-cm Petri dish with beakers)
About 5 mL of 50% H2SO 4 is poured into a 50-mL beaker. The beaker is then placed on a Petri dish to which 30 mL water is added, filling the dish halfway. One of the detection reagents listed may be added at this point. A spatulatip amount (about 0.2 g) of Na2SO3 or NaHSO3 is dropped into the acid, and the smaller reaction beaker is immediately covered with the 150-mL beaker. Within seconds, bubbles of SO2 can be seen near the lip of the larger beaker. The gas may be detected using 10–15 drops of universal indicator, potassium permanganate, or anthocyanin, as described above. Disposal: Pour the acidic solution down the sink with running water. This demonstration may also be performed using household products. Muriatic acid (20% HCl, which is approximately 6 M HCl) can be substituted for the sulfuric acid. The reaction will be slower, taking from 5 to 7 minutes. Rust removal products can be used as a source of sodium bisulfite. The dissolved SO2 may be detected by formation of a colorless complex with anthocyanins taken from red cabbage. Another reaction that produces sulfur dioxide is the addition of concentrated sulfuric acid to sugars to yield carbon and water. This demonstration can also be used when introducing the topic of carbohydrates. It uses the following chemicals and equipment: 2 g sucrose crystals or one sugar cube Concentrated H2SO4; CAUTION: be sure to check that the acid is sufficiently concentrated Universal indicator or 0.002 M KMnO4 or anthocyanin Overhead assembly (10-cm Petri dish with beakers)
In the inner beaker 2 g of sucrose is reacted with 5 mL of concentrated H2SO4. Within 2–3 minutes vigorous bubbling of sulfur dioxide can be seen, along with a brownish color as the sugar degrades to carbon and water. When universal indicator is used as a detection reagent for the SO2 gas, the color changes from green to yellow to orange, then red. This demonstration can also be done with a sugar cube. The cube begins to disintegrate dramatically just before the first bubbles of SO2 are released; however, the reaction takes longer to complete and the bubbling is less vigorous. The other reagents described above, permanganate or anthocyanin, can also be used to detect the SO2. Disposal: Pour the acidic solution down the sink with running water. Place carbon residue in a waste container. Nitrogen Dioxide Nitrogen dioxide is a reddish-brown gas with an irritating odor. The health hazard ratings for nitrogen dioxide are 3 for eye contact, 5 for inhalation, and 3 for skin irritation (7). The solubility of the nitrogen dioxide is about 5–50 g per 100 mL water (7), an amount of gas that dissolves easily in 30 mL of water.
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Detecting Nitrogen Dioxide An acid–base indicator in the water is used to detect the pH change as the NO2 dissolves to form an acidic solution: 2NO2 (aq)
HNO2 + NO 3᎑ + H+
With universal indicator the color changes from green (neutral) to red (acidic) in seconds.
Generating Nitrogen Dioxide The production of nitrogen dioxide can be used when discussing topics such as properties of gases, acids and bases, acid rain, and oxidation-reduction. A colorful demonstration generating nitrogen dioxide is performed by dropping copper metal into 6 M nitric acid. Cu(s) + 2NO3᎑(aq) + 4H + → Cu2+(aq) + 2NO2(g) + 2H 2O(l) The products include brown NO2 and blue Cu II. The following chemicals and equipment are needed: Copper foil (about 0.2 g; a piece 1 cm × 1 cm × 0.02 cm) 6 M HNO3; CAUTION: be sure to check that the acid is sufficiently concentrated Universal indicator Overhead assembly (10-cm Petri dish with beakers)
Put about 5 mL of 6 M nitric acid in a 50-mL beaker and place on a Petri dish. Add water to the Petri dish. The 10–15 drops of indicator should be added later, allowing the color of the CuII to stand out first. Place a second Petri dish with just water and indicator next to this assembly so that the changes in indicator color may be compared later. The copper metal is placed in the acid and covered at once with the larger beaker, to contain the NO2. The 0.2 g of Cu produces about 0.3 g of NO2, which dissolves readily in 30 mL of water. CAUTION: Using too much copper, as for instance the amount in a pre-1982 penny, will produce more NO2 than the water can accommodate, and NO2 will escape into the atmosphere. Within seconds the blue copper(II) can be seen. (Once the blue is obvious, the indicator may be put in the water.) Bubbles of gas appear near the lip of the overturned beaker. After a minute or so the concentration of nitrogen dioxide builds up enough so that the brown color shows, thus ruling out a colorless gas. At this point the color of the indicator is reddish. Disposal: Pour the acidic solution down the sink with running water. Literature Cited 1. Perina, I.; Mihanovic, B. J. Chem. Educ. 1989, 66, 257. 2. CRC Handbook of Laboratory Safety, 2nd ed.; Steere, N. V., Ed.; Chemical Rubber Company: Cleveland, OH, 1971; p 816. 3. Mattson, B.; Anderson, M.; Nguyen, J.; Lannan, J. Chem 13 News 1997, 259, 5. 4. Anthocyanins as Food Colors; Markakis, P., Ed.; Academic: New York, 1982; p 177. 5. Curtright, R.; Rynearson, J.; Markwell, J. J. Chem. Educ. 1996, 73, 306. 6. Salisbury, F.; Ross, C. Plant Physiology; Wadsworth: Belmont, CA, 1985. 7. CRC Handbook of Laboratory Safety, 2nd ed.; Steere, N. V., Ed.; Chemical Rubber Company: Cleveland, OH, 1971; p 794.
Journal of Chemical Education • Vol. 75 No. 12 December 1998 • JChemEd.chem.wisc.edu
