V O L U M E 23, NO. 5, M A Y 1 9 5 1
79
LITERATURE CITED
(1) Assarsuun, G . , Grol Foren. Forh., 63,182 (1941). (2) Barshad, 1.. AN.PL.(‘HEM., 21,1145 (1949). (3) Barshad, I., S o i l S ~ i .66, , 187 (1948). (4) Bertrand, D., C‘onrpt. wrid., 211,406 (1940). ( 5 ) Brockamp, H . , et d . ,A r c h . Lagerstiittenforsch.,77,59 pp. (1944). (6) Ferguson, .I. B., -4m.J . Sei., 37,399 (1914). (7) Ferguson, TV. S..Lewis, A . H.. and Watson, S.J.. I m p . C h r m . Iiids. Jeulolt’s Hi11 Research Sto., B d l . 1 (1940). (8) FBldvari, A , , M a g y a r A411nnii Foldtuni l n t d z e t Ez.i Jeleritdse, Brazdmold. 9, 39 f1947). (9) Goldechmidt. 1.. AI,, h’kvi,fter A-orskr T’idenskaps-dkud., Oslo, I , X a t . S a t u r c . Kluase, 4 (1937). (10) Hevesy, G . yon, arid Hohbie, R . , 2. urzorg. a l l g e w . C‘iieni., 212, 134 11933). (11) Hillebrand. JV. I‘.. S i n . J . Sci.. 6,209 (1898). (12) Hoaglund, P. L., and L m p e , G . A , , A r c h . / z e e v l u r d p h ~ s i u l . , 27,145 (1943). (13) Lakin. H. IT.,Htrrens. T i . E.. and Almond, Hy. Ecoii. Crol., 44, 296 (1949).
(14) Landergren, S., Sveriges Geol. CTndersokn. Ambok, Avhandl. och Uvvsat., 42,S o . 5 (1948). (15) Nickels,- ?M. L., and Rogers,’L. H., IND.ENG.CHEM.,ASAL. ED., 16,137 (1944). (16) Perrin, D. D., .Yew Zealand J . Sci. Technol., 28A,183 (1946). (17) Rankama, K., Bull. comm. ge6l. F i n l a d e , 137,39 pp. (1946). (18) Robinson, W. O., SoilSci., 66, 317 (1948). (19) Robinson, IT. O., U. S.Bureau of Plant Industry, private com-
munication. (20) Sandell, E. E., “Colorimetric Determination of Traces of
Metals,” Sew Tork, Interscience Publishers, 1944. E x . CHEW,ANAL.ED.,28,336 (1936). (21) Sandell, E. B., IND. (22) Sandell, E. B., and Goldich, S.S., J . Geol., 51, 167 (1943). (23) Schneiderhohn, H., rt al., .l-riaes Jtriirb. M i n e r d . Geol., Nonatsh., 1949A,50-72. (24) Westergard, A . H.. Szwiges Geol. L-ndersdhn, y displacrnient of mercury :*rid by displacement of helium. From the difference in denpit,ies, a total pore volume of 0.298 ml. per gram \xis computed which correspoiids to 192 nil. of riitrogm at snturation. Therefore, t l w drsoiption 1)r:tnuh of the isotherm was redetermined using an initial volume of nitrogen n-ell ill PSCPSS of 192 ml. The r ~ s u l is t erhowii in the figure as the “correct desorption isotherm.” From these okxwvations it is upparelit that it is unwise to attempt to determine total pore volumes of adsorbents \vhich popsess a considerable volume in large pores by means of ndsorptiontlcwrption isotherm. I t also appears t h a t the use of an oxygen
’ Prezent addrr.6,
Haiieh a n d Sons Co., Baltimore, Md.
TOTAL PORE VOLUME FROM DENSITY MTA,
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Figtire 1.
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I
/
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!
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~ ~ d ~ r p t i o n - D e s o r p t i oIsotherm Ii for Nitrogen Sample 452 at -195” C .
thermometer to measure the saturation pressure of nitrogen may result in errors of greater magnitude than does the use of a nitrogen thermometer.
ANALYTICAL CHEMISTRY
792 To avoid the uncertainty resulting from the attempt to estimate total pore volumes from adsorption isotherms, Innes suggests that the volume adsorbed a t PIP0 = 0.97 be taken as the measure of total pore volume ( 3 ) . This appears to be an appropriate convention for isotherms of types I, IV, or V, but it is unsuitable for isotherms of types I1 or 111. For the example shown in Figure 1, the volume of nitrogen adsorbed at a relative pressure of 0.97 is 147.5 ml. per gram or 0.229 ml. per gram of liquid nitrogen a t its boiling point. This differs from the total pore volume determined from the density data by 0.069 ml. per gram of liquid or 23.2%. The data also indicate that failure to achieve saturation not only leads to erroneous results a t high relative pressures, but also yields an erroneous desorption isotherm a t all relative pressures above that a t which the desorption branch rejoins the adsorption branch. According to Brunauer ( d ) , desorption from a relative pressure less than unity results in a scanning of the hysteresis loop.
The data of Figure 1 appear to indicate that this is only approximately true. The observations suggest not only that accurate estimations of pore volume cannot be made by means of gas adsorption measurements in the case of type I1 and I11 isotherms, but also that to be certain of obtaining reliable desorption isotherms of these types, a measurement of total pore volume by means of density determinations should be made to provide assurance that saturation of the sample has actually been accomplished prior to desorption. LITERATURE CITED
(1) Brunauer, S., “The Adsorption of Gases and Vapors,” p. 150, Princeton, N. J., Princeton Univer5ity Press. (2)
Ibid.,p. 399.
(3) Innes, W. B., ANAL.CHEM.,23,759 (1951j .
RECEIVEDOctober 12, 1950. Investigation performed anricr multiple fellowship of Baugh a n d Sons Co., Baltimore, 3rd.
Evaporation Errors in Determination of Trace Concentrations of l o w Molecular Weight Solutes in Carbon Tetrachloride ill. R . JIEEKS, 1.. E. WHITTIER. i > D C. U. YOUNG The D o u Chemical Co., .\lidland. .%lich. T H E quantitative analysis of very low concentrations of ImayKpolar solutes in nonpolar solvents, the role of evaporation be seriously underestimated. Loss of solute a t an unexpectedly rapid rate under conditions of sample handling normally considered proper can prove to be a major source of analytical error. An esperience of this sort was encountered by the authors in the course of applying infrared methods for the determination of trace amounts of water and ethyl alcohol in carbon tetrachloride and of water in liquid bromine. I n determinations of tvater in the two solvents a rapid loss or gain occurred a t exposure, depending on the atmospheric humidity. In order to understand these effects and their orders of magnitude, some measurements were carried out on the differential evaporation rates of carbon disulfide, ethyl alcohol, n-butyl alcohol, and acetic acid from solutions of each in carbon tetrachloride. These compounds were chosen to illustrate the effects of boiling point, molecular Tyeight, and polar or nonpolar nature of the solute. Equivalent data on water were not obtained owing to the complication of atmospheric humidity. The significance of these measurements in the field of trace analysis made it seem worth n-hile to present them here.
vapor pressure diagrams. For the binar:. jq‘stems studied, a t the pure carbon tetrachloride end of ~t vapor pressure diagram qualitatively we always have the situation as illustrated by Figure 2. From the usual arguments, it is seen that if a portion of the solut’ionis volatilized, the vapor will be richer i n the minor constituent and the liquid remaining behind will approach pure rarhon tetrachloride in composition.
Y
0
=; Y
9.0
z
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Ia
EXPERIMENTAL
Solutions in carbon tetrachloride of ethyl alcohol, carbon disulfide, and n-butyl alcohol were made up a t concentrations of 3.3 millimoles per mole. Approximately 40 ml. of alcohol solutions were placed in Petri dishes (90 X 20 mm.) and allowed to evaporate from the open dishes a t room temperature. Samples were analyzed for alcohol after various amounts of solution had evaporated. A similar experiment was performed for carbon disulfide in carbon tetrachloride except that, for sampling purposes, 90 ml. of solution were placed in the Petri dishes. The results of the experiments, summarized in Figure 1, show the concentration of solute against per cent loss in weight of solution. The analyses were carried out with a lithium fluoride prism spectrometer. The alcohols were measured a t the 2 . 7 hydroxyl ~ band, and the carbon disulfide a t its 4 . 7 ~band, using the wellknown base-line method ( 1 ) . A cell 3.5 mm. thick was used for the alcohols. For acetic acid, which was investigated less thoroughly, the carbonyl band a t 5.84~was measured using a cell length of 3.5 mm. The determination of carbon disulfide required a 5.0-cm. cell. DISCUSSION
In order to understand the significance of the results indicated by Figure 1, it is well to recall some simple considerations of
1
MW.76
0.5
0
z W
y
0.1
0
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FROM DILUTE SOLUTION IN CCI,
I j
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FIGURE
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0.01
0
10
20
30
40
50
60
70
% OF SOLUTION LOST
Qualitatively the results of the experiments come out as expected. However, Figure 1 shows that after only 12y0 of the carbon tetrachloride solution has evaporated, the concentration of ethyl alcohol falls to 10% of the starting concentration. This rapid loss of solute is a t first glance surprising, perhaps because of a tendency to think of the volatility of ethyl alcohol in terms of the properties of the pure liquid.
