H. D. Gesser, Caroline Lithown, D. Branston, and Ian Thompson
University of Manitoba Winnipeg, Canada
(
I
Capillary T U ~ @Experiments ~ O I Introductory Chemistry Laboratory
The present trend in introductory chemistry t,oward the "physical" approach has been met by the publication of several good texts. The laboratory course has, however, been neglected with the net result that the laboratory does not always compliment the course. One obvious reason is that physical chemistry experiments generally involve elaborate and costly equipment which cannot conveniently be used by large (greater than 200) groups of students. An excellent hypodermic syringe technique (1) has been developed which does not however lend itself t o large classes or to great variety in experiments. It is with the view t o overcome this difficulty that these laboratory experiments were developed.' The requirements rere: inexpensive apparatus which can he issued as locker equipment to each student, and easily mastered teehniques, readily completed in the threehour laboratory period, which give results which compare reasonably well with the literature. The method (2) somewhat similar to that developed by Radley (3) consists of t,he measurement of the volume of expansion at constant Dressure as t,he tem~era&re is changed. his is done by eoufiniug air, the liquid sample, and its vapors in a uniform
If TI is sufficiently low, then: PT - PI = PT.If T211/Tl&= rr, then P. = P T ( ~ a)
(2)
Thus, the partial pressure of a gas at various temperatures can be calculated from a knowledge of the atmospheric pressure, and the sample-mercury distances. Alternatively, the temperature can be maintained constant for kinetic reactions in which a gas is evolved and 1 can be measured as a function of time. This method has been applied to: (1) the vapor pressure of some pure liquids and solids with the values of the heats of vaporization and heats of fusion to he calculated, (2) the gas-solid equilibrium constant of a system with calculations of thermodynamic values, and (3) the kinetics of a reaction with the determination of activation energy. The apparatus is prepared completely by the student who seals the capillary tube (Kimble 3-mm od tubing (approx. 1.5 mm id)) and draws his own eye droppers for filling the tube with liquid sample and mercury. The length of mercury (E) is from 5 to 10 mm, and the mercury to sample distances (l), about 30 mm, is measured with a plastic 6-in. ruler (15 om) available free from scientific supply houses. The ruler should be cut down the middle t o separate the centimeter scale from
mircur; plug. TGS is shown in Figure 1. As the temperature is increased, or if a gas is evolved, the volume expands and the partial pressure of the gas or vapor increases hut the Sample total pressure remains consant since the mercury plug moves t o acFigure 1. Diagram of the capilc o r n d a t e any pressure lo~samplotube. . . change. Let P = the atmospheric pressure exerted on the system, and E = plug length, then PT = P E. It is possible to show that:
+
where PI is the partial pressure of the gas (e.g., vapor pressure) in the system a t T , with 11 the distance between sample (solid or liquid) and the mercury plug. Pa,&, and T1represent a similar set of results.
Figure 2. V ~ p o rpressure of iwpropyt alcohol. 0 results obtoinsd while incremling temperature. 0 results obtained while decreasing tomperatwe. Literature values are represented b y the drawn line.
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the inch scale. I t can then be conveniently bound to the thermometer and capillary tube by rubber bands and mounted on a stand. A 1000 ml beaker initially filled with ice at 0°C serves as the bath for vapor pressures and equilibrium experiments. Both temperature T and length 1 are measured and recorded as the temperature is raised and then lowered between O°C and ahout 50°C with care being taken to allow for equilibrium to be established at each temperature. For a more uniform movement of the mercury plug and to prevent the sample from creeping up the tube, it may be desirable to silanize the glass by treating the tube with a 5% solution of Drifilm (General Electric) in benzene.
Equilibrium
Ammonium Carbamate /O NHnC
'om, decomposes to urea
2
(NHr-C-NH,)
and water (Hz01 a t 130°C, but a t lower tempemtures the solid decomposes reversibly to ammonia and Cot. (4).
Vapor Pressure
The method can he conveniently applied to the determination of the vapor pressure of water, n-CaH70H, and iso-CaH7OH. For t-C4H90H (mp 24"C), the vapor pressure of the solid and the liquid allows the heat of fusion to he determined. A typical plot of the vapor pressure of isopropyl alcohol is shown in Figure 2. The approximation PT- PI = PT,will be invalid when PL2 50 mm (e.g., acetone, methanol). However, a t higher temperatures (when Tz is large) the error introduced becomes less significant and it is possible, by successive approximations, to obtain very good results. This is shown in Figure 3 for acetone.
Figure 3.
Vapor pressure of acetone showing successive approximations.
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Thus, K, = [NH3I2X [Cog]since [NH2-C--ONHI] is solid. It can be shown that
+
where the total pressure P = P N A ~PCO* and P can be calculated from eqn. (2) from the measurements of 1 as a function of T over the temperature range of 0 to 55°C. From the values of P and eqn. (3), K , can be calculated at each temperature. A plot of log K, against 1/T shown in Figure 4 gives AHa and a plot of AG" (or - RT lnKp) against T shown in Figure 5
Figure 4 . A plot of log Kp against 1/T. 0 rervlh obtained while increasing temperature. results obtoined while decreasing temperature. Literature yrrluer ore represented b y the drown line.
measurably slow at 0°C but has a half life of about 7 min between 50 and 55°C. The reaction can be conveniently followed by the evolution of the Nz and it can be shown that the N2formed, (PNJ PNI= (1 - a ) P p - PHIO
I
2X)
I
I
ZM
I
I
290
34)
310
.
I
PO
T°K Figure 5. A plot of AGO against T. 0 results obtained while increasing temperature. 0 results obtoined while dscroaring ternperdure. Litemture values are represented by the drawn line.
(5)
where PH?O is the vapor pressure of HaO at the reaction temperature T* whereas Tlis a much slower temperature, e.g., 0°C. A first order plot of log (P, - PN,) against time as shown in Figure 6 gives the specific rate constant. The benzene diazonium chloride solution is prepared by adding 21 ml of concentrated HC1 to 6.5 ml of pure redistilled aniline and cooling. A cold solution of 4.9 gm NaNO, in 70 ml HzO is added to the cold aniline hydrochloride and the solution diluted to 250 ml with cold water and stored at 0°C or lower. The solution is good for about 3 hours. The cold sample is placed in a treated precooled capillary and the mercury plug added. With T I a t O 0 , 11 is determined. The capillary and ruler are then placed in a 600-ml beaker of water which is maintained at the desired temperature. Readings of 1 are taken a t various time intervals. The rate constant for the reaction can be conveniently determined a t 45, 50, and 55'C. An Arrhenius plot of log k against 1/T is shown in Figure 7. Though the
gives Ma. Students can calculate AH,", AG, and So for ammonium carbamate by completing the following table: AH" kcal/mole
AG' kcalfmole
Soeu!
The ammonium carbamate is added as a powder of about 50 mesh in order to avoid a dead space error. Chemical Kinetics
The reaction
which has been studied extensively (5) is almost im-
Figure 7. An Arrhenius plot of log k again* 1 I T . at 45,50, and 55-C. 0 literature values.
0 rnults obtained
rate constants obtained were about 40% too high, the activation energy 25.0 kcal/mole agreed reasonably well with the literature (5c) value of 28.5 kcal/mole. A possible explanation for the discrepancy in rate constants cannot lie in the solubility of N, in the solution nor to self heating, but possibly to the change in solubility of N2 in solution between OaC and the reaction temperature. tin* lul Figure 6.
A flwt order plot of log (P,
- PNI)
againsttime.
Conclusion
A simple and accurate experiment has been described Volume 44, Number 7, July 1967
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389
which is well suited for the modern introductory course" The technique has been to experiments on vapor pressures, gas-solid equilibrium and chemical kinetics, and could be extended to i1lustrate Raoultk Law 'or other chemical systems.
Complete experimentd procedures (with msignment problems) which have been used in our course by 800 students will be provided interested readers upon request to one of the authors (H. D. G.).
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Journal of Chemical Education
Literature Cited (1) (a) DAVENPORT, D. A,, J. CHEM.EDUC.,39, 252 (1962). (b) DAVENPORT, D. A,, AND SOBA,A. N.. J. CHEM. EDUC., 39.617 (1962). (2) SEE;, C. Y., HERRMMN, R. A., Anal. Cham., 32, 418
AND
E. A., AND JOHNSON, P., Tram. (5) (a) MOELWYN-HUGHES, Far. Soc., 36,954 (1940). (b) CROBSLEY, M. L., KIENLE, R. H., AND BENBROOK, C . H., J. Am. Chem. Soc., 62, 1400 (1940). (c) BORCEARDT, H. J., AND DANIELS, F., J . Am. C h m . Soc., 79.41 (1957).
