In the Laboratory
A Convenient, Inexpensive, and Environmentally Friendly Method of Measuring the Vapor Pressure of a Liquid as a Function of Temperature James H. Burness Department of Chemistry, The Pennsylvania State University, York Campus, 1031 Edgecomb Ave., York, PA 17403-3398 The measurement of the vapor pressure of a liquid as a function of temperature is a standard experiment for laboratory courses from the high school to the university level. The experiment is attractive because it provides a practical application of the Clausius–Clapeyron equation, gives reasonably accurate results, and requires the students to prepare and interpret graphical data. Unfortunately, significant disadvantages sometimes include the use of mercury manometers and the need for large volumes of water to operate aspirators. Van Hecke (1) provides references to a number of experiments, published in this Journal, which are based on the measurement of vapor pressure as a function of temperature. In a significant step, he advocates the use of a vacuum gauge as a replacement for the mercury manometer. Unfortunately, the apparatus requires specialized and expensive pieces of equipment, such as a vacuum pump, needle valves, a quartz immersion heater, and a Variac. Even the replacement for the manometer, the vacuum gauge, costs about one hundred dollars. In a related paper, Berka et al. (2) describe an adaptation of the use of a glass syringe method [based on the Freeman Separate #1163 (3)] to examine the temperature dependence of the vapor pressure of methanol. The authors also compare the results of two methods of obtaining the vapor pressure: calculation using the ideal gas equation and Dalton’s law of partial pressures, and the use of a Chempac pressure sensor interfaced with a computer. In general, the authors found that better results were obtained without the use of the Chempac sensor; on the other hand, the use of the sensor allowed a much wider temperature range. Aside from this dilemma, the experiment is easily performed, inexpensive, and environmentally friendly. Values for the enthalpy of vaporization, calculated from the linear plot of logarithm of vapor pressure versus reciprocal temperature, were within 5% of the actual value in five of eight runs. Predicted boiling points were generally within four degrees (6%) of the actual value. The paper presented here describes a different approach to the measurement of vapor pressures. Like the syringe method, the experiment is inexpensive, environmentally friendly, and easy to set up. Advantages include the direct measurement of vapor pressure from a vacuum gauge, a wide temperature range, and reasonably accurate results. Description of the Modification For years, we had performed the vapor pressure experiment using a simpler (and much less expensive) variation of the isoteniscope method described above (1), with the exception that a water aspirator was substituted for the vacuum pump and a mercury manometer was still used. A sketch of the apparatus is shown in Figure 1. StuPresented at the 13th Biennial Conference on Chemical Education. Bucknell University, July 31 – August 4, 1994.
Figure 1. Sketch of conventional apparatus.
dents first trapped a small bubble of air in the closed end of the bent glass tube (a crude isoteniscope1) containing an organic liquid, then repeatedly grew and fragmented the bubble so that for all practical purposes the bubble consisted almost entirely of organic vapor to allow liquidvapor equilibrium to be established. For each measured temperature (at increments of approximately 10 °C), the pressure was adjusted so that the liquid levels in the two arms of the “isoteniscope” were the same, and then the vapor pressure was read directly from the difference between the mercury levels of the manometer. Students spent two weeks on the experiment; ethanol was used during the first week to provide a “practice run”. Students could work up their data and compare their calculated enthalpy of vaporization and normal boiling point to the actual values for ethanol. For the second week, they were assigned an unknown organic liquid and identified the liquid from a posted list of unknowns and their corresponding enthalpies of vaporization and normal boiling points. Note that a vacuum manifold, with a trap bottle, a ballast bottle, and two stopcocks, was needed to control the pressure in the system. Obvious environmental disadvantages included the presence of the mercury manometers and the usage of large volumes of water during the experiment. Furthermore, the water pressure dropped when a large number of aspirators were being used simultaneously, thereby limiting the extent to which the pressure in the system could be reduced. Aside from the environmental problems, students had trouble with the apparatus. In particular, it sometimes took them a long time to assemble it, and they often (despite repeated warnings by the instructor) allowed large changes in pressure to occur, sometimes producing a “mercury hammer” effect that could have potentially broken the manometer arm. A simple modification to this experiment was made by introducing the use of a hand-held vacuum pump,2 available from most major chemical equipment vendors for about $35. The modified apparatus is depicted in Figure 2. This modification eliminates the need for the mercury manometer and requires no running water, yet
Vol. 73 No. 10 October 1996 • Journal of Chemical Education
967
In the Laboratory
Figure 2. Sketch of modified apparatus.
still allows the students to see changes in pressure (on the built-in gauge) as the temperature is changed. The modified experiment is also much simpler: it takes less than five minutes to set it up, no calibrations are necessary, and there are no vacuum manifold valves or trap bottles. The U-shaped bend in the tygon tubing is designed to trap any liquid that might boil over from the isoteniscope. The remainder of this paper will describe the experimental procedure and present a comparison of the results obtained by both the original and the modified procedures. Experimental Procedure (This section is written as though it were being read by students.) You should work in pairs on this experiment, which lasts two weeks. During the first week you will work with a known liquid (ethanol). In the second week you will receive a small test tube containing approximately 4 mL of a numbered unknown organic liquid. You will collect vapor pressure and temperature data, using the same apparatus and experimental procedure, for each of the liquids. The data from week #1 (known) can be used to check the accuracy and precision of the results, thereby giving some indication of your experimental technique. The data from week #2 (unknown) will be the basis of the lab report that you will submit. You will be provided with a 1-liter beaker, a glass isoteniscope, and the hand-held vacuum pump with its attached tubing. You should assemble the apparatus as shown in Figure 2 using other equipment from your drawer(s). The assembly should be airtight (or almost airtight; if there is a very slow leak, the experiment can still be performed). Connect the isoteniscope to the tubing and test the vacuum release valve by squeezing the pump handle once, then carefully opening the release valve. Now disconnect the isoteniscope and add some organic liquid, using the Pasteur pipette (provided). WARNING: Some organic vapors are poisonous or flammable. Do not leave the liquid in an open vessel for very long. When disposing of organic liquid,
968
pour it into the capped disposal jar (provided). Fill the isoteniscope until the level in the open end is about four centimeters below the level in the closed end. Now it is necessary to trap a small bubble of organic vapor in the closed end of the isoteniscope. (The presence of this bubble ensures that both a vapor and a liquid phase will always be present so that the system can be at equilibrium.) Introduce a small bubble of air into the closed end by carefully tilting the isoteniscope (and using your finger to hold a piece of Parafilm over the open end), then quickly repositioning it after the small bubble of air goes around the bend in the glass. Now, by positioning the isoteniscope so that the closed segment is parallel to the floor (using your hand to hold the open segment above the closed segment), strike or flick the closed section of the tube with your finger, breaking the bubble into many smaller ones. As these form, tilt the isoteniscope to allow all but one tiny bubble to escape into the open end. Your goal is to isolate one very tiny bubble in the closed end. Next, this bubble must be expanded and further fragmented. Reconnect the isoteniscope to your apparatus. Carefully squeeze the hand pump to lower the pressure in the system while you observe the small bubble. As soon as the bubble grows appreciably in size, fragment it as before, retaining only one small bubble. You may notice a number of small bubbles forming at the lower pressure. This could be dissolved air degassing from the liquid. It is important to remove this air (why?); fragment several times if necessary and make sure that the degassing has stopped. At the same time, however, make sure that the pressure is not too low; this could cause the liquid to boil. Warm the closed end of the isoteniscope with your hand to cause the bubble to grow. (Why does it grow?) Fragment as before. The small trapped bubble now consists almost entirely of organic vapor, and it provides the nucleus for the vapor phase that you will produce in this experiment. Fill the liter beaker with cold tap water, and place your thermometer in a clamp so that it is suspended in the middle of the water. It is important that the bulb of the thermometer not be resting on the bottom of the beaker. Likewise, make sure that the bend in the isoteniscope isn’t too close to the bottom of the beaker, and ensure that its closed end is entirely submerged (why?). Lower the pressure in your apparatus slowly, keeping an eye on your trapped vapor bubble. Allow the bubble to grow until the liquid level in the closed arm of the isoteniscope is below the level in the open arm. At this point, the vapor pressure of the trapped bubble is greater than the pressure on the open end of the isoteniscope. (If your pump is unable to lower the pressure enough, warm your water bath by about 10 °C and try again.) Now, very carefully open the relief valve very slightly, allowing air to enter until the liquid levels in the two arms of the isoteniscope are identical. Read the vacuum gauge to the nearest 0.1 inch Hg3 and record the value in your laboratory notebook. Also record the temperature of your water bath as accurately as your thermometer permits. If your system has a slow leak that cannot be corrected, you should grow your vapor bubble as described above. Then simply allow your slow leak to admit air to the system. One partner should watch the isoteniscope and, when he/she sees that the liquid levels are identical, the other partner should read the gauge. Note that the gauge reading is not equal to the vapor pressure. In addition to converting to the proper units, (the vapor pressure should be expressed in torr) you will also need to use the barometric pressure to calculate the vapor pressure of the
Journal of Chemical Education • Vol. 73 No. 10 October 1996
In the Laboratory
liquid from the gauge reading. (How?) After the first measurement, heat your water bath until it is about 8 °C warmer. (But read the temperature as accurately as your thermometer permits.) If the bubble begins growing too large as you heat, carefully bleed in a little air to compress it some more. Now repeat the above procedure to measure the vapor pressure at this temperature. The most likely cause of inadequate data in this experiment is failure to establish true temperature equilibrium at each point. The temperature of the bath will continue to rise after the burner is removed. You should anticipate this behavior and not record the temperature until it has stopped increasing. More important, be sure to allow enough time for the liquid in the isoteniscope to attain the temperature of the bath. Shaking the isoteniscope cautiously to agitate the liquid will hasten this process. Establish the pressure balance at the same time that you read the (stable) temperature. Note that it might be easier to make the levels of liquid in the isoteniscope equal by minor adjustments in the temperature rather than alternately reducing pressure with the vacuum gauge and increasing pressure with the vacuum release valve. Continue warming and taking data at intervals of about 8 °C until you reach a temperature where the vapor pressure exceeds atmospheric pressure (the bubble will grow uncontrollably even though your apparatus is at atmospheric pressure) or until your water bath reaches 80 °C, whichever occurs first. To confirm the equilibrium (reversible) nature of your measurements, now take additional vapor pressure values at lower temperatures between your last and first values. These will thus be within the range of the measurements already made, but will not duplicate them. By adding some cold water to the bath (and removing some of the water in it, if necessary) lower the temperature by 10 to 20 degrees. Adjust the system pressure and determine the vapor pressure of the sample. Repeat this procedure until you have obtained two or three cooling points. (When plotting your data, use a different symbol for the values obtained during the cooling run, to distinguish them from the ones obtained on heating.) You should be able to collect at least twelve data points. BE SURE TO DISPOSE OF YOUR UNKNOWN LIQUID IN THE CAPPED LIQUID WASTE DISPOSAL CAN.
Discussion Whether students performed the experiment using the original or modified apparatus, their data were entered into a spreadsheet immediately after the lab period so they could check the linearity of their plot and see the degree of scatter in the data points. If they were looking at the results of the first week’s work, they could also see their experimentally determined values for the enthalpy of vaporization and normal boiling point of ethanol, and the corresponding percentage errors. What the students did not know, however, was that after they had viewed their results, their raw data were captured by use of a spreadsheet macro. In this manner, it was possible to determine percentage errors, correlation coefficients, etc. for all students without having to worry about mistakes they might have made in working up the results. This information was obviously useful for grading the laboratory report, but an additional benefit was that the capture of these data permitted comparison of the results for the two types of apparatus. (A copy of this Lotus 1-2-3 compatible template is available from the author if a diskette is supplied). A comparison was made between a group of nine pairs of students who had performed the experiment using the conventional apparatus and another group of eight pairs
Table 1. List of Unknown Liquids Liquid 2-monotone
Enthalpy of Vaporization Normal Boiling Point (kJ/mole) (°C) 22.2
222
cyclohexane
32.8
80.7
dioxane
35.8
101
ethanol
40.5
78.3
t-butyl alcohol
42.1
82.5
n-butyl alcohol
45.9
117
49.9
161
65-95 (depending on age of sample)
130-247
cyclohexanol isogeritol
Treatment of Data Draw two graphs based on your data from the second week: (1) vapor pressure vs. temperature and (2) natural logarithm of vapor pressure vs. reciprocal absolute temperature. Perform a linear least-squares regression to determine the best straight-line fit to the data points. From the slope and intercept of the best straight line for the second plot, calculate the enthalpy of vaporization for the liquid and the constant C in the equation, ln P = (∆HV/R) (1/T) + C Your instructor will have posted a list of organic compounds with their heats of vaporization and normal boiling points (see Table 1). Determine the boiling point of your unknown by two methods: (1) use your value of ∆H V and C, and the Clausius–Clapeyron equation, to calculate the normal boiling point; (2) determine the boiling point graphically using your linear plot (remember that the normal boiling point is defined as the temperature at which the vapor pressure of the liquid is 760 torr). Compare your values of ∆HV and the normal boiling point with the tabulated values and identify your compound from the list.4
Figure 3. Plot of ln Pv vs. reciprocal temperature. Typical student results for t -butyl alcohol using the hand pump. The correlation coefficient refers to the data regression line.
Vol. 73 No. 10 October 1996 • Journal of Chemical Education
969
In the Laboratory
of students who had run the experiment using the handheld vacuum pump. A typical plot, based on actual student results for t-butyl alcohol using the modified apparatus, is shown in Figure 3. The literature regression line was calculated by using the average of the enthalpies of vaporization at 25 °C and at the normal boiling point (4). The enthalpy of vaporization and the normal boiling point calculated from the student-obtained results agree to within 3% and 2%, respectively, of the actual values. It is important to note that the hand pump modification allows students to collect many data points during the lab period (typically, only about 8 to 10 points could be collected using the original apparatus). In fact, the number of points seems to be more limited by the suggested increment of 8 °C than by the time available; many students finish early or have ample time to repeat the experiment if they make a major error. The correlation coefficient from the least-squares regression fit is a direct measure of the consistency of the students’ results. Figure 4 shows a line graph of these quantities. This graph indicates that the modified apparatus produces results that are at least as consistent (and perhaps somewhat more so) than the original experimental setup. In any case, this figure shows that the experiment allows a high degree of reproducibility. The percentage errors could be determined for all students by comparing their results with the known values for the samples they used. These errors are a measure of the accuracy of their results. There were two pairs of students in each group with unusually high percentage errors (from 15% to 30%). It is significant that these students had high percentage errors for both the ∆Hv and the boiling point determinations, strongly indicating faulty technique. If these outlying values are discarded, the average percentage errors shown in Figure 5 are obtained. It is interesting that these errors are exactly those reported by Berka et al. for the syringe method. Since the typical error is on the order of 5%, we do not use all of the unknowns shown in the Table; rather, we use unknowns that have reasonable differences in their enthalpies of vaporization and boiling points. For example, cyclohexanol, t-butyl alcohol, and n-butyl alcohol are good choices for unknowns; they can be easily differentiated once both the enthalpy of vaporization and normal boiling points have been determined, and they are environmentally safe. Conclusion The use of a hand-held vacuum pump to replace both a water aspirator and a mercury manometer allows students to determine the vapor pressure of a liquid as a function of temperature easily, quickly, and accurately, and the experiment is environmentally friendly. This modification permits the use of a wide temperature range and makes it possible to collect a large number of data points during the experiment. Acknowledgments The author is very grateful to Roger Twitchell for helpful discussions concerning both the experimental modification and this manuscript. Thanks are also extended to one of the author’s students, Tom Grab, for drawing Figures 1 and 2. Notes 1. The isoteniscope is made by sealing one end of a 30-cm
970
Figure 4. Comparison of least-squares correlation coefficients.
Figure 5. Comparison of percentage errors. C = conventional apparatus; HP = modified apparatus with hand-held pump.
length of 7-mm fling glass tubing and putting a 150° bend in it at approximately 10 cm from the sealed end. 2. Mityvac Superpump D , repairable, 35 mL (item #6210); Neward Enterprises, Inc. P. O. Box 725, Rancho Cucamonga, CA 91729-0725. 3. Although the inner scale of the gauge is calibrated in units of cm Hg, the triangular needle is narrower at the outer scale, and more accurate estimates of the pressure can be made if the outer scale (in in. Hg) is used. 4. A referee has suggested that some instructors, in lieu of using unknown liquids, may want to tell their students which compound they have been assigned. The students can then took up the literature values and calculate the percentage error for their determination.
Literature Cited 1. Van Hecke, G. R. J. Chem. Educ. 1992, 69, 681 and references cited therein. 2. Berka, L. H.; Kildahl, N. K.; Bergin, S. J.; Burns, D. S. J. Chem. Educ. 1994, 71, 441. 3. Hagen, J. W. “Equilibrium Vapor Pressure of a Liquid as a Function of Temperature”, Laboratory Studies in General Chemistry; W. H. Freeman: San Francisco, 1972; p 1163. 4. Lide, D.R., Editor-in-Chief. CRC Handbook of Chemistry and Physics, 71st ed.; CRC: Boca Raton, FL, 1990.
Journal of Chemical Education • Vol. 73 No. 10 October 1996
