In the Laboratory
The Determination of the Percent of Oxygen in Air Using a Gas Pressure Sensor
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James Gordon* and Katherine Chancey Division of Science and Mathematics, Central Methodist University, Fayette, MO, 65248; *
[email protected] As seen from a number of publications (1–6), the determination of the percent of molecular oxygen in the air continues to be a favorite experiment in introductory chemistry courses. The experiments are relatively fast and efficient at determining the quantity of what is perhaps the most critical gas in the atmosphere. Several authors suggest using steel wool in their experimental design (1–5). In these experiments, the iron in the steel wool reacts with gaseous oxygen from the air to form rust as represented in the simplified reaction below: 4Fe(s) + 3O2(g)
2Fe2O3(s)
The reaction is carried out in a closed container so that the finite quantity of oxygen can be measured by reacting it with the iron in the steel wool. Several recipes can be found in the references above for the type of steel wool and the preparation of the steel wool using various concentrations of acetic acid or vinegar. The acid serves as a catalyst for the reaction. Most of the methods in the references involve measuring the height of a water column or mass or volume of water collected as a result of the decrease in pressure in the test chamber as the steel wool reacts with molecular oxygen. In this work, methods in which water is used to help determine the percent of oxygen will be referred to as water-measurement methods. The work presented here differs from the water-measurement methods in that a direct change in pressure of the system was recorded. A Vernier LabPro data collection system, similar to the Texas Instruments Calculator-Based Laboratory system (TI-CBL), connected to a TI-83 calculator and interfaced to a gas pressure sensor was used to determine the decrease in the total gas pressure inside an enclosed test tube as steel wool reacted with the gaseous oxygen. This experiment is performed in a general chemistry lab in which students compare the results calculated from the pressure measurements obtained with the calculator-based systems to those obtained in a water-measurement method similar to that outlined by Postma et al. (3). It has been observed that students gain experience with the calculator-based data collection systems, enjoy using them, and appreciate the ability to collect real-time data. This experiment is suitable for any college or high school introductory chemistry course.
Pieces of coarse #3-size steel wool, weighing approximately 1.0 g, were cleaned by soaking them in acetone for 30 seconds. The steel wool was thoroughly dried after removal from the acetone and then soaked in 50 mL of white vinegar (about 0.8 M acetic acid). After drying, the steel wool was soaked in a diluted vinegar solution (about 0.08 M acetic acid). The steel wool was thoroughly dried, spread out, and
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Hazards Acetone is flammable and should not be used around any ignition source. There are very limited hazards associated with the use of steel wool and vinegar. The steel wool can cause eye irritation and upper respiratory disease. It can also generate flammable hydrogen gas when in contact with mineral acids. The vinegar can cause irritation to the nose, throat, and lungs. Results and Conclusions
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quickly placed into a 125-mm test tube. Lastly, the test tube assembly was connected to the pressure sensor using a onehole 00-size rubber stopper for a snug fit. If the steel wool is not thoroughly dried of the acid there is a possibility that hydrogen gas will be generated by the reaction of the steel wool with the excess acid rather the desired reaction with gaseous oxygen. It was found that even “thorough” drying with paper towels leaves enough residual acid to catalyze the reaction. A Vernier LabPro data collection system controlled by a TI-83 Plus calculator running the Datamate software was set up to collect data in the “time-graph” mode. Data were collected for 30 minutes at a rate of 4 points per minute. The test tube and sensor assembly was left undisturbed until the data collection was finished. The percent difference between the initial atmospheric pressure and ending steady-state pressure was calculated and used as the percent of oxygen. The same procedure as outlined above was used to prepare the steel wool for determining the percent of oxygen in air for a water-measurement method. Once the steel wool was in the test tube, the test tube was inverted and placed into a 250-mL beaker of water. The water level inside the test tube was checked every five minutes until it stopped rising, which indicated that the oxygen supply inside the tube had been depleted. The volume of water inside the test tube was carefully recovered and measured using a graduated cylinder. The steel wool was removed from the test tube and the total volume of the test tube was measured. The percent of oxygen was determined by dividing the volume of the water accumulated as a result of the reaction by the total volume of the test tube.
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A typical graph of gas pressure versus time is shown in Figure 1. The data collection was started approximately one minute before the test tube was connected to the pressure sensor so that the initial pressure could be recorded. Following the connection of the test tube to the sensor system a dramatic drop in the pressure was observed until the final steady state pressure was reached, indicating the oxygen had been depleted. The average value for the percent of oxygen
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In the Laboratory
Pressure / atm
1.00 0.95 0.90 0.85 0.80 0.75 0
5
10
15
20
25
30
35
Time / min Figure 1. A typical pressure versus time curve for the reaction between oxygen and steel wool.
drogen gas is generated by a side reaction of the iron with the acid catalyst. A third possibility is that the increased pressure was caused by the presence of water vapor that can be seen from the condensation of water inside the test tube as the reaction proceeds. Lastly, perhaps the small fluctuations in the gas pressure observed after 30 minutes indicate that even with the catalyst present the last quantities of oxygen take a longer time to react. It should be noted that this discrepancy in the measurement is likely present in the watermeasurement methods. However, it is masked by the large standard deviation. In conclusion, this experiment allows students to explore a fundamental reaction used in the determination of oxygen. The use of current technology allows them to collect reproducible data and to compare the results to those of a classical experimental procedure. W
in the air was found to be 19.4 ± 0.4% (n = 7). This value was compared to the average value of 20.6 ± 1.4% (n =8) from the water-measurement method. The percent oxygen for the calculator-based pressure method is more precise than the water-measurement method. However, the percent of oxygen determined by the pressure method was somewhat lower than expected. Several reasons have been considered for the discrepancy. Since the reaction is exothermic, perhaps the increased temperature leads to a small but significant increase in the pressure inside the test tube. In order to obtain a final percent of oxygen of 21% from the data in Figure 1, it would be necessary for the final steady-state pressure to have been 0.776 atm, which is 0.015 atm lower than the observed steady-state value. Using an alternative form of the Charles and Gay–Lussac relationship of measuring pressure as a function of temperature with volume and moles of the gas remaining constant, the final temperature would have had to increase from 25 ⬚C to 32 ⬚C to give rise to such an increase in pressure. The test tube was felt before and after the several experimental runs with no detection by touch of such a large rise in temperature. A second possibility is that a small but significant quantity of hy-
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Supplemental Material
Instructions for the students and notes for the instructor are available in this issue of JCE Online. Acknowledgments The authors would like to thank Alan James for helpful discussions during the revisions of this manuscript. We would also like to thank Central Methodist University for funding this research project. Literature Cited 1. Gettys, N. S.; Jacobsen, E. K. J. Chem. Educ. 2001, 78, 512A. 2. Braathen, P. C. J. Chem. Educ. 2000, 77, 1410. 3. Postma, J. M.; Roberts, J. L.; Hollenberg, J. L. Chemistry in the Laboratory, 4th ed.; W. H. Freeman: New York, 1997; pp 201–207. 4. Birk, J. P.; McGrath, L.; Gunter, S. K. J. Chem. Educ. 1981, 58, 804–805. 5. Martins, G. F. J. Chem. Educ. 1987, 64, 809–810. 6. Najdoski, M.; Petrusevski, V. M. J. Chem. Educ. 2000, 77, 1447–1448.
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