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 and 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
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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
I
FROM DILUTE SOLUTION IN CCI,
I j
I
FIGURE
I
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.
V O L U M E 23, NO. 5, M A Y 1 9 5 1
793
In thtlse very dilute solutions, the solute behavior is expressed by Heni,y’s law: P, = S,k where k is a constant chaiacteristic of the solvent-solute system. The Henry’s law constant, in thermodynamic terms, is an expression of the escaping tendency, or fugacity, of the solute. This ~ high for a substance like ethyl alcohol, because may b ( aLilormally the assoc,iation effect from hj-droxyl bonding is eliminated in very dilute solution We are then dealing with a low molecular weight d solvent-solute species, : l ~ ~ the interactiim twcomes one common w to pol~r-nonpolar systems gena r~raliy. Even a t.uhstance like n-butyl a alcohol tends to escam from cara bon tetrachloride, although the R effect is rather slight. I t is interesting fcs note the difference of x X’ 3
The determination of water in nonpolar solvents has been found to be subject to all the effects mentioned, to an extreme degree. It might have been more to the point to illustrate this discussion with a water-containing system. But the kind of experiment performed for the systems chosen would offer formidable difficulties when water is involved. Under ordinary conditions the experimenter would be operating in an atmosphere containing amounts of solute (water) significant with respect to the solution concentrations of interest. Evaporating carbon tetrachloride containing small amounts of water leads to either a decrease or an increase in water content. An equilibrium is reached with atmospheric water vapor concentration depending on the humidity. If these observations are accepted, several rather obvious precautions must be taken in handling solutions of this kind for analysis. Determinations on hundreds of samples analyzed independently by both chemical and spectroscopic means gave very erratic results until the sources of error were successfully traced back to handling procedures ordinarily considered entirely adequate. Samples must be kept hermetically sealed. Transferring a sample from one vessel to another must be done with considerable care. Pouring a carbon tetrachloride solution containing a few parts per million of water, in such a way as to expose a considerable surface of solution to the atmosphere, can greatly affect the concentration of water. Simple calculation may show that the amount of vapor space above the solution in a container may be significant with respect to the solution volume itself. The investigator must be prepared to ask himself if the determination of water content has any meaning. According to the history of the sample, water content may be more a function of atmospheric conditions coupled with sample handling than anything else. Only when the sample container is handled with very special care will water concentration be anything more than an expression of solvent-atmospheric water vapor equilibrium.
rg +
99.93 1007. behavior tretneen carbon disulfide CCI, ~] \vliich h a w and ~ L - ~ J U I ,alcohol. very n e a r l y e q u a l n i v l e c u l a r Figure 2 weights, but are v e different ~ with respect to polar rharacter Acetic acid was also iiivestigated to some estent. At 0.1% concentration very little change in concentration is noted upon partially evaporating the solution. But a t lower concentrations, about O.Olyo,there if h marked l o ~ e r i n gof acetic acid content with evaporation. This i.q probably explained by the fact that a t ~ is present to a high degree as the the higher concentration r l i acid dinirir. However, a t 0.01% and loner, there is a noteworthy incr(mc of monomer to dinier ratio as shown by the infrared spectrum, P O t h a t a t very low concentrations a low molecular weight species i k prrwnt (60 for the monomer as against 120 for the dimer). .is a i.ule. the authors concluded that the loss of solute is proportionately greater the loner the concentration for the conipounds investigated. If this latter observation is correct, the neces.4tj- for emphasis on extreme care in sample handling is increased at minute concentrations for solutions such as those considercd lirrr
LITERATURE CITED (1) Wright,
s..ISD. ESG.C H E Y . , A4N.%L.E D . .
13, 1 (1941).
RECEIVED June 5 , 1950. From a paper presented a t the Pittsburgh Conference on Analytical Chemistry and Applied Spectroscopy, Joint Meeting, Division of Analytical Chemistry, Pittsburgh Section. ~karERIcaNCHEMICAL SOCIETY,and Spectroscopy Society of Pittsburgh, February 13 t o 17, 1950.
Spectrographic Determination of Silicon in Uranyl Nitrate Solutions F. T. BIRKS .4tomic Energy Research Establishment, Harwell, Didcot, Berkshire, England I L I C O S has been determined in uranous uranic oxide (LT308) and in uranium metal and compounds after conversion to U308, according t o the direct current arc carrier distillation method ( 1 ) . and in uranium metal by Kalsh (b),who used graphite counter e!ectrodw in the controlled alternating current arc. A number of uranyl nitrate solutions, each containing about 300 mg. of uranium per nil. of dilute nitric acid solution, were to be examined for silicon. The carrier distillation method was considered too slow for this purpose, as it would involve the conversion of the whole or most of the sample t o U308 and the subsequent operations of mixing with the carrier are time-consuming. Because the only impurity element sought was silicon and there were sensitive silicon lines located in a region free from an excessive number of uranium lines, it was decided to apply a direct hurn procedure and to use uranium itself as an internal standard PREPARATION OF STANDARDS
A pure sample of U308was prepared by extraction of strong uranyl nitrate solution with ether, removal of the ether under reduced pressure, and final ignition of the residue to u308. Pure precipitated silica was diluted with this base by thorough grinding in an agate mortar to obtain the standards: 5000, 1000, 500,
100, 50, and 10 p.p.m. of silicon relative to UsOs. A blank estimation of silicon was carried out on the pure U308 used for the dilution. PRELIMINARY INVESTIGATION
The electrodes were shaped from National Carbon Co. pure graphite rod 0.5 inch (1.25 cm.) in diameter to fit on a graphite support 0.125 inch in diameter, and had a crater 6 / 3 2 inch in diameter X 1/20 inch deep. They xere preburned for 30 seconds a t 10 amperes before loading, to remove surface contamination and reduce the silicon in the electrode material to a uniform low level. A moving plate exposure was first made on a 6-mg. charge in a 10-ampere direct current arc to study the relative emission rates of silicon and uranium. It was found that the silicon intensities followed the uranium fairly closely and fell off rapidly after 20 seconds becoming zero a t 30 seconds. The exposure conditions adopted were 30 seconds a t 10 amperes with the sample as anode. To investigate the effect of charge weight on the result, weights of 4,5,6,and 7 mg. of the 500 p.p.m. standard were exposed; the intensity ratio was found to remain reasonably constant. It had been shown in connection with other work t h a t in general the sensitivity of trace elements added in the form of an aqueous solution was higher if the solution were prevented from soaking into the porous graphite electrode by the presence of a thin grease film. Furthermore, nonpenetration of the electrode
