In the Laboratory
Determination of Enthalpy of Vaporization Using a Microwave Oven A. P. Kennedy, Sr. Department of Chemistry, North Carolina A&T State University, Greensboro, NC 27411
The typical liquid–vapor phase experiment involves measuring either the vapor pressure of the liquid under constant temperature conditions or the boiling point under constant pressure conditions (1–4). Both experiments are normally performed at or below atmospheric pressures. No undergraduate experiments reported in the literature have measured the vapor pressure at elevated temperatures and pressures. With the recent advancements in research-grade microwave ovens it is now possible to safely measure the vapor pressure and temperature of polar liquids at elevated temperatures and pressures. A simple undergraduate laboratory experiment utilizing this technology is described below.
Table 1. Enthalpy of Vaporization Experimentally Determined and Calculated from CRC Values (5 ) Liquid
∆H Literature (cal/mol)
∆H Experimental (cal/mol)
% Error
Methanol
8617.9
8350.9
3.09
Ethanol
9154.1
8721.7
4.95
Acetone
7147.4
7395.8
{3.35
Isopropanol
9397.3
9155.5
2.64
A research-grade CEM MES 2000 microwave oven was used in these experiments. It consisted of a test vessel containing a fiber optic temperature sensor and a pressure sensor. The oven had the additional safety feature of automatically shutting off if any liquid vapors leaked from the test vessel. The test vessel was made of microwave-invisible polymeric materials. The inner liquid lining was made of Teflon; the outer containment sheath was made of Ultem. Methanol, ethanol, acetone, and isopropanol were used in these experiments. Chlorinated liquids such as methylene chloride should not be used because of their reactivity with the Ultem sheath. The test vessel was sealed and the experiment was performed under constant volume conditions. Each run required approximately 40 min. A run consisted of microwave heating for 10 min and a 30-min thermal cool down. The initial microwave heating was very rapid and produced a hysteresis in the heating-vs.-cooling curves (see Fig. 1). This occurred because of the rapid nonequilib-
rium heating, which caused a temperature gradient between the liquid and the gas phase. This hysteresis is not observed when the optical temperature probe is placed in the gas phase. The data collected during thermal cooling was used to calculate the thermodynamic data, since they were under near-equilibrium conditions. The pressure and temperature data collected during a single run provided all the data necessary to construct the corresponding liquid– vapor phase curve. The data were collected via a RS232 interface and imported into a DOS graphing package. Twenty-five–milliliter samples were placed in the 100mL test vessel and exposed to microwave powers between 300 and 400 W. Either the maximum pressure (200 psi) or the maximum temperature (200 °C) can be used as the controlling variable. For these different polar liquids a maximum pressure of 180 psi was sufficient to collect the data and still be well below the specified safety limits of the microwave oven. The measured pressure was corrected to account for the partial vapor pressure of air that was initially present in the test cell. The resulting Clausius–Clapeyron relationship for methanol gave excellent agreement with literature
Figure 1. Hysteresis of temperature vs. pressure response for fiber optic probe immersed in liquid and in gas phase. Each symbol represents one minute of data collection.
Figure 2. Clausius–Clapeyron plot comparing experimental data from microwave oven and data from CRC (5 ).
Experimental Procedure
Vol. 74 No. 10 October 1997 • Journal of Chemical Education
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In the Laboratory values (5) (see Fig. 2). The agreement between literature and experiment for some common liquids is shown in Table 1. Deviations from high-pressure data are less than 5%. This simple, safe experimental procedure can be applied to a variety of polar liquids that are microwave active. Unlike conventional vapor phase equilibrium experiments, it provides the student with further insight into how the pressure and temperature of an equilibrium system change simultaneously. Acknowledgment
Literature Cited 1. Shoemaker, D. P.; Garland, Carl W.; Nibler, J. W. Experiments in Physical Chemistry, 5th ed.; McGraw-Hill: New York, 1989; p 219. 2. Tobey, S. W. J. Chem. Educ. 1962, 39, 258. 3. Berka, L. H.; Kildahl, S. J.; Bergin, S. J.; Burns, D. S., J. Chem. Educ. 1994, 71, 441. 4. Burns, S. J.; Berka, L. H.; Kildahl, S. J. J. Chem. Educ. 1993, 70, A100. 5. Handbook of Chemistry and Physics, 59th ed.; Lide, D. R. Ed.; CRC: Boston, 1978–1979; pp D240–D257.
I would like to thank Lois Jassie of CEM Corporation for the loan of the microwave oven.
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