In the Laboratory
An Inquiry-Based Chemistry Laboratory Promoting Student Discovery of Gas Laws
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A. M. R. P. Bopegedera Department of Chemistry, Lab I, The Evergreen State College, Olympia, WA 98505;
[email protected] Gas laws are covered in most undergraduate general chemistry courses and even in some high school chemistry courses. Once the concept of pressure and its units are introduced, chemistry texts launch into a discussion of gas laws (1). Experiments to enforce the understanding of gas laws are often done in the laboratory after these concepts are covered in lecture. From a description of Boyle’s original work (2) to simple demonstrations and laboratory experiments (3– 27), many articles on the topic of gas laws have been published in this Journal. An interesting article on the assessment of students’ and teachers’ understanding of gas laws was published recently (28). A laboratory experiment book for high school and college general chemistry, published by Vernier, covers some of the gas laws (29). For the past few years, I have taken a different approach to teaching gas laws by letting students “discover” them in the laboratory using their own lab data. The subsequent lecture time is used for problem solving using gas laws. This pedagogy was effectively used with chemistry majors and nonmajors. I conduct these experiments with readily available, reasonably priced, Vernier software and laboratory equipment.1 This equipment is user friendly and training time is minimal. The hardware and memory capabilities of a typical desktop computer are sufficient to host the software.1 Prior to conducting the lab experiments on gas laws, students are exposed to the following concepts. • Chemical foundations, atoms, molecules, ions
1. Pressure and volume (temperature and amount of gas held constant) 2. Pressure and temperature (volume and amount of gas held constant) 3. Volume and temperature (pressure and amount of gas held constant) 4. Pressure and the number of moles (volume and temperature held constant) to determine the universal gas constant
Equipment and Chemicals The materials needed for these laboratory experiments are listed below. • An inexpensive 30-mL plastic syringe, available at any drugstore, was used as the sample cell for Experiment 1. • For Experiments 2 and 4, a 250-mL Erlenmeyer flask served as the sample cell.
• Introduction to the periodic table • Nomenclature of compounds • The mole concept, molar mass, writing balanced chemical equations • Stoichiometric calculations, limiting reagents • Types of chemical reactions • Acid–base reactions
At The Evergreen State College, students taking general chemistry register for lecture and laboratory classes simultaneously. Therefore it is possible to use either the lecture class or the laboratory to introduce any concept. (For descriptions of the academic environment at The Evergreen State College please see refs 30–33). Four experiments are used to investigate the relationships between properties of gases and only one set of lab equipment was used for each experiment. Since data acquisition time is short with the Vernier equipment, 25 students were able to complete all four experiments in a lab period of three hours. Laboratory time was dedicated to collecting data, which was then exported to Microsoft Excel spreadsheets for graphing and analysis. (Excel is used for data analysis in all the labs throughout the academic year; using it for these labs as well supports students’ skill development. Logger Pro soft-
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ware provided by Vernier could be used also, if desired.) Students analyzed their data individually, outside the lab period. They were directed to plot graphsW and make inferences on the relationships between properties of gases based on these graphs, prior to attending a discussion session to go over their lab work. Each student submitted his or her lab report a few days after this discussion. Students conducted experiments and explored the relationships between the following properties of gases.
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• For Experiment 3, a 50-mL glass syringe was used (Becton-Dickinson and Co., product # 512135). • Three-way valves were purchased from Cole-Parmer (plastic three-way stopcocks with Luer connection, product # C-30600-02). • Atmospheric air was used as the “gas sample” in Experiments 1 and 3. • In addition to atmospheric air, He, CO2, and N2 gases (Matheson Gas, 99.9% purity) were used in Experiment 2. • Students used CO2 gas in Experiment 4. • The instructor repeated this experiment with SF6 and Ne gases (Matheson Gas, 99.9% purity) and provided the results to students for inclusion in their data analysis.
Conducting the Laboratory Work
Experiment 1 The relationship between pressure and volume of a gas was explored while keeping the temperature and the number of moles of gas as constants. The values for the volume
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of the gas were read from the 30-mL plastic syringe and the pressure was recorded with Vernier sensors and software. Students plotted graphs of pressure values versus volume and pressure values versus reciprocal volume and easily inferred that the pressure of a gas is inversely proportional to its volume at constant temperature (Figure 1), and is directly proportional to the inverse of its volume (Figure 2), thus “deriving” Boyle’s law.
Experiment 2
Figure 1. Graph of pressure data values vs volume of a gas sample at constant temperature. Students analyzed these data and could deduce that the pressure of a gas is inversely proportional to its volume at constant temperature.
A fixed amount of gas was held in a constant volume sample cell (250-mL Erlenmeyer flask) and the relationship between pressure and temperature was explored. The experiment was done with four different gases (atmospheric air, He, CO2, and N2). The sample cell was immersed in a water bath at approximately 0 °C and the temperature of the bath was slowly increased by heating. The temperature and corresponding pressure of the gas sample were recorded continuously until the final temperature reached 90 °C. By plotting graphs of pressure versus temperature for the four gases, students concluded that the pressure of a gas sample held at constant volume is proportional to its temperature (Figure 3).
Experiment 3
Figure 2. Graph of pressure data values vs reciprocal volume of a gas sample at constant temperature. After analyzing these data students could deduce that the pressure of a gas is directly proportional to the inverse of its volume at constant temperature.
The relationship between volume and temperature of a gas was explored while keeping the pressure and the number of moles constant. A volume of atmospheric air was drawn into the 50-mL glass syringe, which served as the sample cell. The cell was sealed, immersed in a constant temperature water bath at 0 °C and the inside pressure was allowed to equilibrate to atmospheric pressure. The value for the volume of the gas was read from the syringe and the temperature was recorded. This was repeated at several different temperatures of the water bath. Students plotted a graph of volume versus temperature and deduced that the volume of a gas at constant pressure is proportional to its temperature (Charles’s law).
Experiment 4
Figure 3. Graph of the pressure data values of the CO2 gas sample versus its temperature at constant volume. With these data students could deduce that the pressure of a gas sample held at constant volume is proportional to its temperature.
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The relationship between the amount of gas and its pressure was explored while keeping the volume and the temperature of the gas constant. A 250-mL Erlenmeyer flask (which served as the sample cell) was evacuated and weighed with an analytical balance. A small amount (∼150 torr) of CO2 gas was introduced into the cell, its pressure was recorded and the sample cell was weighed to determine the mass of CO2. This step was repeated several times, each time with a different amount of CO2 in the cell. The volume of the cell was determined by filling the sample cell with water, determining the mass of the water and using the density of water to arrive at the volume of water (which is equal to the volume of the cell). Students input their data into a “class spreadsheet” and these combined data were used in the analysis. The instructor repeated this experiment with Ne and SF6 gases and provided these data to students to be included in their data analysis.W Students were directed to plot graphs of pressure (P) versus the number of moles (n) (which showed
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In the Laboratory
that P is proportional to n at constant V and T ) and PV versus nT and compare the slopes of the PV versus nT graphs for the three different gases, CO2, Ne, and SF6 (see Figures 4 and 5). The data in Table 1 show that Experiment 4 provides remarkably good values for the universal gas constant. Using three different gases made it possible to establish the “universality” of the gas constant. During the discussion following lab work, the different units of the universal gas constant were examined (1). It was a good moment to revisit unit conversion, emphasize its importance and reinforce this skill. Boyle’s law and Charles’s law have been studied previously with Vernier equipment (29). However, in this laboratory exercise—in addition to the unique pedagogical approach of teaching gas laws using the discovery approach in the laboratory—students investigated all the different physical properties of gases and their relationships. Another highlight of this lab experience is that students determined the universal gas constant using laboratory data collected by the entire class. By comparing the value of the gas constant calculated using data from different gases, students were able to recognize that this constant is truly universal.
Figure 4. Graph of pressure data values versus number of moles for CO2, Ne, and SF6 gases at constant volume and temperature.
Hazards Students were permitted to use the gas-handling manifold in Experiment 4 only in the presence of the instructor. Care must be taken not to vent CO2 gas into the laboratory. A CO2 monitor was available to ensure that any leaks would be quickly detected. Instructor’s Observations Year after year, I am pleasantly surprised to note that even at the end of the academic year, students are able to arrive at the gas laws by recalling the graphs they plotted in this laboratory. They also had a better conceptual understanding of the universal gas constant, compared with former students (who did not learn gas laws using this discovery approach) who treated it as a number that must be “looked up” from a table of physical constants, without fully understanding how it is derived or why it is “universal”. From this experience I conclude that students find it easier to remember gas laws when they are derived with students’ own laboratory data instead of memorizing them from a textbook. This laboratory provided an excellent opportunity for collaborative work. Students who had more experience with the software helped those with less experience. Errors made during data collection affected data analyses for the entire class, which helped students recognize more strongly the need for accuracy in the laboratory. Some students had been exposed to gas laws prior to taking this class, either in high school or in a previous chemistry course. I certainly do not discourage students from reading the textbook ahead of lab time, and thereby learning gas laws prior to discovering them in the laboratory. Yet, there is a difference between learning gas laws from a textbook and proving that the gas laws do hold using their own lab data. Gas laws, when discovered in the laboratory, are no longer an abstract concept that must be memorized. In my view,
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Figure 5. Graph of PV values versus nT values for CO2, Ne, and SF6 gases at constant volume and temperature (determination of the universal gas constant).
Table 1. Comparison of Results of Determining the Universal Gas Constant from the Graph of PV versus nT, by Gas Gas
Universal Gas Constant a (atm L K᎑1 mol᎑1)
CO2
0.101
Average Student Experimental Error b (%) 23.1
Ne
0.0823
0.293
SF6
0.0809
1.41
a These values were obtained from the slopes of the PV versus nT graphs shown in Figure 5. The reported value for the universal gas constant is 0.08206 atm L K᎑1 mol᎑1. b Note that the error is highest with the CO2 gas student data, partly because of students’ limited lab experience and also because each student group used a different analytical balance. The error is lowest with Ne gas, the most likely of the gases tested to act as an ideal gas.
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this helps students retain knowledge of gas laws better and longer. The data provided in the Supplemental MaterialW for the four experiments2 could be used to conduct a “dry lab” on gas laws. Although this will somewhat diminish the laboratory discovery experience, students would nevertheless be able to arrive at the same conclusions from graphing these data. W
Supplemental Material
Notes 1. Information on Vernier LabPro software, sensors, specifications for the hardware and memory requirements for the computer are available from the Vernier Products Web site at http:// www.vernier.com (accessed Nov 2006). Similar products may be available from other vendors. 2. Experiment 4 may be done without SF6 and Ne gases. However, including these two gases helped students understand the “universality” of the universal gas constant. Students were not given access to SF6 and Ne gases because of the higher cost of these gases and time constraints.
Literature Cited
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1. Zumdahl, Steven S.; Zumdahl, Susan A. Chemistry, 6th ed.; Houghton Mifflin Company: New York, 2003; also other similar general chemistry textbooks. 2. Neville, R. G. J. Chem. Educ. 1962, 39, 356–359. 3. Lewis, L. D. J. Chem. Educ. 1997, 74, 209–210. 4. Mills, P.; Sweeney, W. V.; Marino, R.; Clarkson, S. J. Chem. Educ. 2000, 77, 1161–1165.
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10. 11. 12. 13.
Laboratory instructions, block diagrams for experimental setups, and student data and analysis for each of the four experiments are available in this issue of JCE Online.
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