NOTES
June, 1962
1197
WATER-RICH EQUILIBRIA I N THE SYSTEM CH&OONa-CH&OOH-HzO1 BY LEONARD C. LABOWITZ~ Department oi Chemistry, New York University, Fashington Spuare Center, New York 8,N . Y . Received November $0,1961
I n 1910 it ‘was reported by Ball@ that solid solutions are formed at both the water-rich and acidrich ends of the phase diagrams in a number of binary mixtures of the type wRCOOH-H~O (where n-RCOOH = HCOOH, CH&OOH, CPH6COOH, and n-C3H,COOH). The purpose of his work was to support the hypothesis popular a t that time* that the deviation of such systems from the simple Raoult-van’t Hoff freezing point depression law was due not to molecular association but to solid solution formation. Particular attention was given by Ball6 to the system acetic acid-water, since it had been the most extensively studied member of the series. I n view of the crystallographic differences between acetic acid and ice, it would seem rather improbable that such solid solutions could exist to any appreciable concentrations. Ordinary ice is hexagonal5 and acetic acid is orthorhombic.6 On the other hand, it is not inconceivable that these differences could be overshadowed by the influence of hydrogen bonding. Considering the current widespread interest in the existence of solid solutions of ice in which ice is the major component,7.s it is surprising that this early work of Ball6 evidently has been unnoticed by recent workers. The purpose of the present investigation is to examine the claim made by Ball6 that solid solutions of ice are formed in the water-rich region of the system acetic acid-water. The original study by Ball6 was based upon chemical analysis of the phases and use of the tracer method of van Biljert.@*laIn the present work, Schreinemakers’ method of wet residues11p12 is applied to the system sodium acetate-acetic acidwater a t -10.2 and -13.9’ in order to determine whether the composition of the solid phase in the water-rich region of the system is pure ice or the reported solid solution. Experimental Apparatus and Procedure.-The apparatus used for the separation of the liquid phase from the wet residue (Fig. I ) ronsisted or a long, narrow test-tube fitted with a fritted glass filter stick (F) and a thermocouple well (C). The (1) This work was performed at Kew York University under the auspices of a National Science Foundation Summer Research Program for College Chemistry Teachers. (2)Department of Chemistry, The City College of New York, New York 31, New York. (3) R. Ba116, 2. physik. Chem., 73, 439 (1910). (4) A. Findlay, A. N. Campbell, and N. 0. Smith, “The Phase Rule and Its Applioations,” 9th Ed., Dover Publications, Inc., New York, N. Y.,p. 156. (5) K.Loqaidale, Proe. Roy. Soo. (London), 8246, 424 (1958). (6) R. E.Jones and I).H. Templeton, Acta Cryst., 11, 484 (1958). (7) L. Pauline, “The Nature of the Chemical Bond,” 3rd Ed., Corne11 University Press, Ithaea, N. Y., 1960, p. 464. (8) L. C. Liabowitz and E. F. Westrum, Jr., J . Phye. Chem., 66, 408 (1961). (9) A. van Biljert, Z . physik. Chew., 8 , 343 (1891). (IO) J. E. R.icci, “The Phase Rule and Heterogeneous Equilibrium,” D. Van Nostrsnd Go., Inn., New York,N. Y., 1951,p. 324. (11) F. A. If. Schreinemakers, 2. physik. Chem., 11. 7 5 (1893). (12) J. E.Ricci. ref. 10, p. 323.
A-
I
B-.
C
- 1 G
H Fig. L-Wet
residues apparatus.
filter stick was suspended in the test-tube through a short section of glass tubing (E), the point of entry being closrd off from the outside atmosphere by means of a short piece of tightly fitting rubber tubing (D). The test-tube was partially filled with sample (about 10 g., completely liquid a t room temperature), and the entire assembly shown in the drawing was immersed for periods of 8 to 16 hr. in a low temperature bath t o the level indicated by (B). The thermocouple wires were shielded from the bath fluid by plastic tubing ( A ) . In the -10.2” study, the low temperature bath consisted of well stirred eutectic brine mixture dewar flask. I n the 13.9”study, a methanol-containing low temperature bath (Tenney Engineering, Inc., Union, N. J., Model No. 7020) capable of maintaining a give12 temperature to within 3 ~ 0 . 2over ~ the range +70 to -20 was used. During the solidification process, the filter stick was kept in the lowered position as shown. To prevent premature entry of material into the filter stick, the tip was closed by a small glass plug attached with rubber tubing (G). Temperature inside the sample container was monitored nith a calibrated single-junction copper-constantan thermocouple and in the bath fluid with a calibrated mercury-in-glass thermometer (in the case of the -10.2” work) and a calibrated pentane-in-glass thermometer (in the case of the -13.9” work). After the attainment of equilibrium, the rubber tubing a t (D) was disconnected from the wider tube, and the filter tube was raised until the plug could be disengaged from the tip by the tube (E). The filter stick then was quickly lowered once again to the bottom of the test-tube and as much liquid as possible -’as withd r a m through the filter stick by suction. To determine whether equilibrium had been established under the conditions used in these experiments, duplicate samples were subjected t o various refrigeration times and starting temperatures (above and below the desired final temperature). The data obtained are self consistent and within the limits of experimental error of the results of other workers on the water solubilities of acetic acid and sodium acetate at - 1 O . Y Analysis and Materials.-Acetic acid was determined by titration using a calibrated buret of a weighed sample with standardized 0.1 N NaOH, using phenophthalein as the indicator. T o determine sodium acetate, a separate samplc was passed through an ion-exchange resin ( Amberlite IR-120, Rohm and Haaai Company), thereby converting the
-
1198
NOTES
Vol. 66
phase in the water-rich region of the system is, within experimental error, pure ice and not the solid solution claimed by Ballb. Aside from mme obvious mistakes in the phase diagrams shown by him,it is believed that the results obtained by Ball6 are incorrect because of misapplication of the method of tracers, misinterpretation of the data, and poiasible errors made by him in the chemical analysis of the barium acetate tracer.
ti110
TABLE I WET RESIDUESDATAFOR THE SYSTEMCH3COOSaCHaCOOH-HI0 Wt. % CHaWt. % COON& CHaCOOH
Tieline no.
1 2 Fig. 2.--\Tlater-rich region of the sgsttin CH8COONaCH3COOH-H20 a t -10.2
.
3
4
5 6
Liquid Residue Liquid Residue Liquid Residue Liquid Residue Liquid Residue Liquid Residue
Mathematically extrapolated terminus
-10.2 f 0.1" 8.98 13.97 0.57 wt. 5.73 9.12 4.49 21.99 .29 wt. 3.51 16.90 12.61 6.83 .34 wt. 9.29 5.12 -13.9 f 0.2" 5.58 27.06 .55 wt. 4.06 18.90 13.10 13.57 .20 wt. 8.87 9.12 .1i wt. 17.11 6.57 11.04 4.31
% CHaCOOH % CH3C001\;a
70 CHSCOOH % CH3COOKa % CHaCOO?;a % CH$OOH
Acknowledgment.-The author wishes to express his sincere appreciation to Profs. John E. Trance and Seymour Z. Lewin of New York University for the generous use of their facilities. The financial support of the National Science Foundation is gratefully acknowledged.
Fig. 3.-Water-rich region of the system CH3COONaCHaCOOH-HzO at - 13.9". salt to the free acid, and total acetic acid was titrated as described above. The sodium acetate then was estimated by the difference. The method was found t o be dependable even in the presence of excess acetic acid. A standard sodium acetate-acetic acid solution was prepared by mixing measured amounts of standardized sodium hydroxide and acetic acid solutions. The results obtained by chemical analysis agreed within 4 parts per thousand of the value predicted on the basis of the concentrations and amounts of the mixed starting materials. I n a different experiment, two separate aliquots of a given solution of sodium acetate (excess acetic acid absent, this time) were analyzed by this method and the results agreed n-ithin 5 parts per thousand of each other. The relative mean deviation for the sodium acetate determination therefore is estimated to be 2.5 parts per thousand. h calibrated buret was employed for all titrations. Commercial reagent grade chemicals without further purification were used throughout.
Results and Discussion According to Ba116, the composition of the saturated solid solution at - 10.2 and - 13.9' would be approximately 14 and 18 weight % CH3COOH, respectively. The results obtained in the present investigation, given in Table I and Fig. 2 and 3, iiidir.:ite, I ~ U ' C T W ,th:ii thc cwmposition of thc solid
THERMODYNANIC STUDIES OF T H E IODINE COMPLEXES OF s-TRITHIAKE, THIACYCLOHEXANE, AND THIACYCLOPENTANE IN CARBON TETRACHLORIDE SOLUTIOK BY J. D. hfCcULLOUGH .k>D
IRlliEL.4
c. ZIMYERMANN
T h e Department of Chemzstry o f the Unsverszty o f CalefoEolnza at Lo8
Angeles, Los Angeles 24, Calzfornza Recesved December l S 8 1961
The present work is an extension of previously reported studies1-3 on systems of the type D - 1 2 = D Iz,where the donor, D, is an organoselenium compound or an organic sulfide.
+
Experimental Materials.-Thiacyclopentane and thiacyclohexane were from the same samples used in ref. 1 and were supplied bq the American Petroleum Institute Research Project 48A, Bartlesville, Oklahoma. s-Trithiane was kindly supplied by Professor E. E. Campaigne of Indiana University. The recrystallized solid melted a t 220-221'. The iodine and carbon tetrachloride and the experimental procedures were those described in ref. 1 and 2. Method of Calculation.-The method of calculation was (1) J. D. McCullough and D. hiuhey. J . A n . Chem. Soc., 81, 1291 11959). 12) J. D. McCiillorlgh and I. C. Zimmermann, J . Phys. Chem., 64, 1064 (1960). (3) 1 I ) \ I I (7iilloiipli n n d T C Z ~ ~ r ~ i n i ~ i r n iul tni dn. ,, 66, XXS (lQOlJ,
