R. J. Kokes, M. K. Dorfman, and T. Mathia
The Johns Hopkins University Baltimore, Maryland
Experiments for general chemistry 11
A Simple Vacuum System
In the freshman chemistry course a t the Johns Hopkins University, a group of experiments are performed with a simple vacuum system. These studies of both homogeneous and heterogeneous systems include the application of Boyle's law; the verification of the ideal gas law; the determination of the The new freshman laboratory course rtt Johns Hopkins has been described in TAIS JOURNAL,39, 16 (1962). Many of the experiments are innovations in an introductory course. This series of urticles describes the experimental procedures in some detail.
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molecular weights of gases; effusion studies; and the application of the gas laws to the vapor in equilibrium with solids, liquids, and solutions. The Apparatus
The simple vacuum system is shown in Figure 1; it consists of a manifold of capillary tubing, and includes a manometer and a spherical bulb which is permanently attached to the capillary manifold. Two standard taper joints are connected to the manifold, one via a stopcock. Auxiliary pieces of equipment include two
spherical bulbs, one of which is equipped with a twoway stopcock, and a cylindrical hulb with a thermocouple well. All three bulbs are equipped with standard taper joints to permit attachment to the manifold. The manifold is connected to the vacuum pump by way of a three-way stopcock; by proper manipulation of this stopcock, the system can he evacuated or opened to the atmosphere. In addition to these items, each apparatus includes a Variac-controlled oven; a thermocouple; and an effusion "plate" consisting of a sealedoff piece of tubing with a pinhole, which can he inserted in the line between the pump and the system. The apparatus is clamped to a rack, the base of which is easily accommodated on a student laboratory bench. Boyle's Law and the Ideal Gas Law
The volumes of the individual parts of their vacuum system are determined by each pair of students by an application of Boyle's law. First, they calculate the volume of one of the bulbs from the weight of water it can hold; then, they dry the hulb, fill it with air a t atmospheric pressure, and replace i t on the evacuated system. By allowing the air to expand into each part of the system in succession and reading the pressure at each expansion, they obtain enough information to calculate the volume of each part from the relation P V = constant a t constant temperature and moles of air. Thc students next test the validity of the ideal gas law applied to air. To do this, they first determine the volume of their cylindrical bulb, described above, then raise the oven to enclose all of the bulb except the stem (which is assumed to he a t room temperature) and increase the temperature in stages regulated roughly by the srttings of a Variac, over a range from room temperat,urc to about 350°C. For each point, the students
read the pressure and determine the temperature from the thermocouple potential, measured with a portable potentiometer. In their report on the results of this experiment, the students compare the observed pressures with those calculated from the ideal gas law, and thus directly demonstrate the behavior of a real gas compared with the "ideal" gas. The Molecular Weights of Gases
The gas laws are applied to determine the molecular weights of several gases by two methods. In the first method, the ideal gas law is directly applied; the students fill a hulb with a measured weight of gas, and allow the gas to expand into the vacuum system. From the known volume and temperature and the measured pressure and weight of the gas, they calculate the molecular weight directly. Typical results of this experiment appear in Table 1. The second method of determining the molecular weight of a gas involves an application of the kinetic theory of gases. In 1846, Thomas Graham showed that the rate of effusion of a gas through a small aperture at constant temperature varies inversely as the square root of the molecular weight. By an argument based on kinetic theory, one can formulate this law as,
where A is the area of the aperture, V the volume of the container, and Po and P Iare the pressures a t the start of the experiment ( t = 0) and at the time, t, after the start of the experiment. The other symbols are standard. In this experiment, a small orifice is attached to the vacuum system as described above. A bulh containing the gas is attached to the evacuated system and some of the gas is bled out until a pressure of about 400 mm is attained. The students then open the stopcock and, as the gas escapes through the orifice, they record simultaneously the pressure and elapsed time. From plots of log P I versus t , the molecular weights of several unknown gases are obtained. The molecular weights obtained with both methods are given in Table 1 below; some typical data are represented graphically in Figure 2. Table 1 Gas Air Butane Helium Areon
True moleculm weight
M.W.
obtained, Ideal Gas Law
M.W. obtained, Graham's Law
29.4 59.2 4.4 40.2
48.6 5.1 35.3
29.2 58.12 4.003 39.9
...
Vapor Pressure Experiments
Vapor pressure experiments are carried out with benzene, which melts at 5.5' C. About 10 ml of the liquid is placed in a spherical bulh with a standard taper joint, connected t,o the manifold via a stopcock; the bulb is surrounded by an ice-salt bath, and the whole system is evacuated until the henzene is free from air. If the benzene remains a super-cooled liquid a t the initial temperature, -10' C, a piece of dry ice is Volume 39, Number 1, January 1962
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touched to the bulb to initiate freezing. The vapor pressure of henzene from -10" C to room temperature is then measured a t intervals of about three degrees, each point being taken when the temperature and pressure reach constant values.
Table 2.
Student Results
Experimental value
Literature value
Melting point (' C) Boiling point (' C) & C) Kb( C) AHvLp (call (cal) Vap. press. a t 20' C (mm HG)
1 0 78.5 9 3.5 6980 1989 76.2
5.53 80.10 5.12 2.57 7353 2351 74.1
Melting point (' C) Boiling point (' C) Vap. press. a t 20' C (mm Hg) Vrto. Dress. a t 20' C, ealc. from Raoult's Lad. (mm Hg)
-6 83.5 65.3 67.6
65.8
Rcnzm~ ~
(o
Conclusion
,
0.71 0
10
20
30
40
I 'SO
60
70
Time (recondsl Figvre 2.
Student rervltr in effusion experiment.
The vapor pressure of a solut.ion of naphthalene in benzene is also studied, the only modification in procedure being a correction for the change in concentration due to the loss of some benzene during the out gassing. The students calculate the true concentrat,ion from the weights of the solution before and after the experiment. The students are told to interpret their data in the light of the ClausiusClapeyron equation and Raoult's law. They use the graphical representation? shown in Figure 3, to obtain the melting and boiling points of the pure liquid and the solution; the melting point was taken as the temperature at which the cnrves for the solid and liquid intersect, and the boiling point. is estimated by extrapolating the curve for the liquid to atmospheric pressure. With this information and the concentration of the naphthalene-benzene solution, the students calculate the molal freezing point depression and boiling point elevation for benzene. Knowing the form of the relation between vapor pressure and temperature, they use the slopes of their curves to calculate the heats of fusion and vaporization of benzene; and finally, they test Raoult's law by comparing the observed and calculated vapor pressures of the solution a t 20" C. The data and results obtained by a high school senior, who helped test the experiment, appear in Fignre 3 and Table 2. For comparison, values of the vapor pressures of benzene a t various temperatures taken from the literature ( 1 ) have been treated in the same way as the students' results.
The students acquire the elementary vacuum techniques with little damage to the apparatus. Most important, in these experiments, the students study these ideal gas law relations in systems in which the properties are directly measured and unobscured by side effects. The students discover not only the varioua aspects of a dynamic equilibrium but also the fact that this equilibrium is not instantaneously attained. (More often than not in the course of this experiment
Figvre 3. The vapor prerwre of benzene: e represents student dot.; solid lines are from lnternotionoi Critical Tables.
the benzene supercools.) Thermodynamic quantities are made tangible when they 'alculate enthalpies of fusion and vaporization from the familiar quantities, pressure and temperature; the relations governing colligative properties are found to be a simple consequence of Raoult's law coupled with the ClausiusClapeyron equation. Thus, in one experiment., the students learn from their own experience that chemistry is a set of integrated principles which lose their vigor when compartmentalized.
EDITOR'S NOTE: Additional experiments in this program are described in papers which will appear in the February 1962 issue of THIS JOURNAL.
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