In the Laboratory
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Unknown Gases: Student-Designed Experiments in the Introductory Laboratory John Hanson* and Tim Hoyt Department of Chemistry, University of Puget Sound, Tacoma, WA 98416; *
[email protected] Designing an experiment is an intellectually exciting and creative activity, and successfully implementing an experiment you have designed is one of the most rewarding aspects of science. Many upper-level science students, particularly those engaged in research, have the opportunity to participate in this important part of the scientific process, but relatively few laboratory activities for introductory students provide them with the chance to actually design an experiment, probably owing to the difficulties associated with finding appropriate problems that they can successfully solve given their modest chemical knowledge and limited exposure to laboratory techniques. For the past nine years we have included a laboratory experience for our introductory students in which they design and implement an experiment to determine the identity of three unknown gases. We do not give them instructions on how to accomplish this; we simply pose the following problem: You will work on your own in this experiment. As part of the lab this week you will, individually, be assigned three unknown gases from the following list: air, nitrogen, oxygen, argon, carbon dioxide, helium, methane, hydrogen.1 You must design experiments to determine the identity of your three gases. (HINT: Look at your lecture notes.) Your TA will dispense the gases to you in any containers you wish. You will be given one sample of each of your gases. In addition, you may request one sample of any gas by name or request another sample of any one of your unknowns.
Students are instructed that for safety reasons they may not inhale any of the gases and any use of flames in the lab must be approved by the instructor. In addition, students are told that when they come to the laboratory their notebooks should contain a written description of their experimental plans as well as the molecular structure, molar mass, and relevant reactivity information for each of the gases. This is not a trivial problem for the students to solve— these are all colorless, odorless gases! 2 But the identity of the gases can be determined without the use of any specialized equipment or techniques, and the conceptual background students need can be obtained from the discussion of gases presented in most introductory chemistry texts. Permitting students to request one additional sample of any gas by name (or another sample of any one of their unknowns) plays an important role in letting them double check their methods or observe the properties of a known sample. Mass- and Density-Related Observations Differences in the molar masses of the gases are reflected in differences in the densities of the gases and in the masses of a given volume of the gas (Table 1). Even a qualitative comparison of gas densities or masses can be used to assign an unknown gas to one of 4 groups: hydrogen and helium are much less dense than air (a balloon containing them will
rise); methane is significantly less dense than air (a small methane balloon will usually not float, but is clearly “lighter” than a balloon containing air); nitrogen, air, and oxygen are similar to each other; and argon and carbon dioxide are clearly “heavier” than the other gases. Many introductory students, having just learned the ideal gas law, are eager to take a quantitative approach by weighing their unknown gases and using this information (along with the pressure, temperature, and volume) to determine the molar masses of them. In fact, this is the basis for several lab experiments and lecture demonstrations described in this Journal (1–3). From the values in Table 1 it is clear that a balance of 0.01-g accuracy can be used to differentiate between the gases, even when relatively small amounts (100 mL) are used. Many students are initially surprised when they find that the weight of their container filled with one of their unknown gases is actually less than without it! (Assuming that they haven’t used a pump to evacuate the container ahead of time.) This is a chance for some very productive discussions about the buoyant force of air (4 ). Approximate volumes, pressures, and temperatures are accurate enough to distinguish many of the gases, but most students actually go to the effort to measure them. We have a thermometer and barometer set up in the lab and glassware is available for measuring volumes. There are many opportunities during the lab to discuss with students the relative merits of various experimental strategies, the accuracy needed for reliably distinguishing between different gases, or the validity of their assumptions. For example, the pressure in a fixed-volume container filled with a gas and then stoppered is close to atmospheric pressure, but what is the pressure in a balloon? In a Mylar-type balloon or a plastic bag the pressure is close to atmospheric, but the pressure of a gas in a latex rubber balloon will be higher than that of the atmosphere owing to the tendency of the rubber to
Table 1. Properties of Gases Gas
Molar Density/ ∆Mass/ Reaction with Burning Splint gb Mass/g (g/L) a
Hydrogen
2.01
0.084
᎑1.120
Gas rapidly combusts ("pops")
Helium
4.00
0.166
᎑1.038
Extinguishes splint
16.03
0.667
᎑0.537
Gas combusts (flame is seen) Extinguishes splint
Methane Nitrogen
28.02
1.165
᎑0.039
Dry air
28.97
1.204
0.000
Oxygen
32.00
1.331
+0.127
Splint burns more rapidly
Argon
39.95
1.661
+0.457
Extinguishes splint
Carbon dioxide 44.01
1.830
+0.626
Extinguishes splint
Splint burns normally
aCalculated
using the formula d = P mmol/RT ( P = 1 atm; mmol is the molar mass; R = 0.0821 L atm/mol K; T = 293 K). bMass of l L of gas minus mass of 1 L of air. This is what will be observed when the mass of a container (filled with air) is subtracted from the mass of the same container filled with the gas.
JChemEd.chem.wisc.edu • Vol. 79 No. 7 July 2002 • Journal of Chemical Education
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In the Laboratory
contract. If this issue gets raised, students can readily measure the increased internal pressure of the gas in a balloon by using a U-tube manometer (3). (A maximally filled latex balloon has an internal pressure about 20 torr above atmospheric, so even if students don’t correct for this it doesn’t significantly affect their results.) Alternatively, students can measure the mass of a balloon containing a known gas and then calculate the pressure. Reactivity (Combustion) Observations Students usually realize that there will be obvious reactivity differences observed if the gases are exposed to a flame (Table 1): hydrogen and methane are both flammable, air supports combustion, oxygen speeds up combustion, and the other gases extinguish the flame. In our experience a fixedvolume container of 250 mL or less (e.g., an Erlenmeyer flask) that has been flushed with hydrogen or methane may be safely ignited by removing the top and carefully bringing a burning splint up to the edge of the opening. Under these conditions hydrogen produces an audible pop as it rapidly combusts, and methane produces a small flame that continues burning for a few moments after the splint is removed. For additional safety, instructors may wish to have these combustion tests done behind a small blast shield. Balloons filled with one liter or less of hydrogen or methane (about the size of a cantaloupe) may be safely ignited by suspending them in the air with a string and then igniting them with a burning splint that is attached to the end of a stick at least one meter in length (5, 6 ). The balloon should be at least 2 feet from all surfaces and the flame should be brought up to the far side of the balloon, away from the person igniting it. Mylar balloons should not be ignited (5). There are other tests that students may have devised while doing research for the lab. We try to accommodate them if possible. For example, one useful test is the detection of carbon dioxide using limewater (7). Hazards Hydrogen and methane are highly flammable, so any use of flames in the lab should be regulated by the instructor. Combustion tests should be carefully supervised and performed only on limited amounts of the pure gases, never on mixtures. Students are not permitted to inhale any of the gases; consequently any observation that requires inhalation may not be used. Even when students do not purposely inhale the gases, it is inevitable that small amounts of these gases will be released into the room and consequently we have selected only gases that are nontoxic. It is also recommended
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that a well-ventilated laboratory room be used for this experiment. Conclusions This laboratory is an extremely positive experience for a wide variety of introductory students and is often mentioned as the favorite lab from the first-semester course. Students become very interested and intrigued by the challenge of designing their own experiment. Unlike some labs where a detailed experimental procedure is provided, this lab forces students to think through both the theoretical and the practical aspects of the lab beforehand. (It was extremely rewarding to walk through the school cafeteria one day and hear a group of students in an animated discussion about various strategies for determining the identity of their unknown gases.) Although the problem is challenging, nearly all students are able to identify their unknown gases and they feel a real sense of pride and accomplishment when they do. W
Supplemental Material
Additional notes to instructors on the logistics of performing this laboratory are available in this issue of JCE Online. Notes 1. These gases were chosen for their low toxicity and ready availability. We have recently added propane as an additional unknown gas. 2. The house natural gas that we use as the source of our methane contains a low concentration of odorant to permit detection of leaks. This does not significantly change the apparent mass of the gas. Students are informed that they may not inhale the gases, so they should not use any information about the odor in their analysis.
Literature Cited 1. Jackson, B. A.; Crouse, D. J. J. Chem. Educ. 1998, 75, 997– 998. Lieu, V. T.; Kalbus, G. E. J. Chem. Educ. 2002, 79, 473. 2. Bodner, G. M.; Magginnis, L. J. J. Chem. Educ. 1985, 62, 434–435. 3. Zaborowski, L. M. J. Chem. Educ. 1972, 49, 361. 4. Ball, D. W. J. Chem. Educ. 1998, 75, 726. 5. Battino, R.; Battino, B. S.; Scharlin, P. J. Chem. Educ. 1992, 69, 921–923. 6. Corkern, W. H.; Hughes, E. Jr. J. Chem. Educ. 1999, 76, 794. 7. VanCleave, J. P. Chemistry for Every Kid: 101 Easy Experiments That Really Work; Wiley: New York, 1997; pp 70–73.
Journal of Chemical Education • Vol. 79 No. 7 July 2002 • JChemEd.chem.wisc.edu
