753
NOTES
June, 1958
TABLE I Temp.,
OK.
XAI
0.036 ,055 .088
.lo1 ,128 .186 .188 .330 .340 .461, .544 .549 .650
83.0
70.0 In P A ,
-0.0223 - .0428 .3764 .5371 .6461 ,9735 .9462 1.2698 1.2223 1.3712 1.4782 1.4584 1.5059
XAr
0.060 .087 .139 ,165 ,276 .381
0.9610 1.3455 1.7435 2.0497 2.2029 2.3052
0.293 .375 .388 .495 .540 .570 .615
In P A *
3.0963 3.2638 3.2577 3.3725 3.4388 3.4662 3.4904
XAr
89.2 In
0.009 ,019 ,024 .049 .052 .054 .066 .096 .221 .353 .463 .581
96.2 XAr
PAr
1.4386 2.1268 2.1242 2.2757 2.4027 2.8783 2.9210 3.2107 3.6354 3.8841 4.0539 4.1685
In PA,
2,7559 3.2232 3.7762 3.6205 3.9235 3,9402 4.0920 4.1449 4.1965
0.058 .080
.loo
.117 .136 ,171 .183 ,209 .221
TABLE I1 Argon at 83OK.
Krypton at 105’K.
Xenon
Ref.
a L. (cal./mole) 1840 2550 Heat of sublimation b -8.3 K(cma2/kg.X lo5) -9.9 Isothermal compressibility C 7.6 Cp(cal./deg.) 8.0 Heat capacity (constnnt pressure) d v(cm.8) 28 32 Molar volume 3.61 4.05 e 3.42 Low temp. collision diameter 180 230 e E/k 124 Equilibrium pair energy . (”) .. “Selected Values of Chemical Thermodynamic Properties,” Circular 500, Series I, National Bureau of Standards, U. Govt. Printing Office, p. 544-545. b The value for argon represents an extrapolation of the data while that for krypton was selected because V / K -9000 cal. which was thought to be order of magnitude correct; J . W. Stewart, Phys. .Rev., 97, 578 (1955). C K. Clusius, 2. physik. Chem., B31, 459 (1936). d L‘Smithsonian Physical Tables,” Vol. 88, Smithsonian Institution, Washington, D. C., 1934, p. 159. e J. 0. Hirschfelder, R. B. Bird and E. L. Spotz, Chem. Revs.,4 4 , 205 (1949).
4.)
s.
for the random-mixed model of the strictly regular solution was evaluated as before. The value WAB = 320 - 1.7T cal. per mole, with an uncertainty of 50 cal. was obtained. This value produces the solid line in Fig. 1. The dashed line is the best fit obtained with a temperature-independent WAB. The extrapolated critical temperature for phase separation is 56 i 8°K. The large uncertainty in this estimate is caused by the rather long extrapolation. Discussion of Results.-Both the entropy and energy parts of WAB are one-half the corresponding values for the krypton-xenon system. We have attempted to calculate these quantities according t o several current theories, but have been hampered by the lack of good data on the pure solids, especially for xenon. We have used the “cell model” of Prigogine and Bellemans,s the “refined average potential model” of these authors and Mathot,’ and the “one liquid” and “two liquid” theories of Scott.8 The assembled data used, and their sources are presented in Table 11. The results of the computations are given in Table 111, in terms of the excess thermodynamic quantities at IL: = 0.5. If one bears in mind that all these theories are full of approximations, and that the data used are often rather poor, it is clear that the theories give a reasonable account of the krypton-xenon system. The heat term for the argon-krypton system is fairly well predicted, but all the theories fail to account for the excess entropy for this system. On (6) I. Prigogine and A. Bellemans, Disc. Faraday Soc., 16, 80 (1953). (7) I. Prigogine, A. Bellemans and V. Mathot, “The Molecular Theory of Solutions,” North Holland Publishing Co.,Amsterdam, 1957, Chapt. X. ( 8 ) R. L. Scott, J . Chem. P h y s . , 25, 193 (1956).
the other hand, as we have noted above, the simple “corresponding states” idea that the ratios of heat and entropy terms for the two systems should be the same, gives a good account of the s y ~ t e m . ~ TABLE I11 Soln. argon-krypton Temp. = 83OK. X = 0.5
Measured “Cell model” “One liquid” (‘TWO liquid” “Refined av. potential model”
Excess free energy of mixing 771
Exoesa enthalpy
of mixing FI‘
(cal.jmole) (oal.j;nole)
Excess entropy of mixing S‘ (e.11.)
0.43 .14 .06 .03
45 30 41 27
42 46 30
41
31
81 53
170 113
0.85 .57
139
149
.12
80
-
.12
Soln. krypton-xenon Temp. = 105OK. X = 0.5
Measured “Cell model” “Refined av. potential model”
(9) G. D.Halsey and M. P. Freeman, Nature, 178, 431 (1950).
THE SOLUBILITY OF CALCIUM SOAPS BY JOHNTHOMAS YOKE,I11 Miami Valley Laboratories, Procler and Gamble Co., Cincinnati, Ohio Received January 6 , 1968
It has seemed worthwhile to redetermine the solubilities of calcium stearate, palmitate, laurate and oleate in water. The results of previous direct determinations1-4 are in very poor agreement. (1) J. Zink and R . Liere, 2. angew. Chem., 28,
[I] 229 (1915).
754
NOTES
Vol. 62
The result of an indirect determination5 of the solubility of calcium stearate disagrees with the values based on analysis of the saturated solution. In the direct determinations it is in all cases uncertain that equilibrium was attained, and there has been little concern with purity of the samples, or with the nature of the solid phase involved. In several cases the experimental errors have exceeded the solubility values obtained. A radio-tracer technique involving Ca45was used for analysis of the saturated solutions.
If a given portion of solid phase is equilibrated with water, the resulting solution decaiited and discarded, and the process repeated, the (more soluble) impurities should be preferentially leached out of the solid phase. After a number of leaching cycles, and after sufficient time of agitation with a new aqueous phase to re-establish equilibrium, a true solubility value should be obtained, which would not change with further agitation or leaching. This was indeed found to be the case in this work. All final solubility values were obtained in this Preparation of the Calcium Soaps.-Fatty acids of purity manner with one exception. The exception is calcium oleate at 60", where degreater than 99yo(by setting point) were used as the starting materials.6 The theoretical amount of NaOH was composition of the sample (to a brown oil) occurred added to an acid in hot 95% ethanol. The radioactive within 200 hours, despite such precautions as calcium soap was precipitated by addition of the theoretical amount of a mixture of ordinary reagent CaCln with a Ca46 purging the system with purified nitrogen and (Oak Ridge) enriched CaClz sample (ca. 1 millicurie per shielding the system from light. Therefore, magram of calcium soap). The precipitate was filtered, dried terial for this determination was pre-leached a t a and then placed in enough boiling anhydrous n-butanol to lower temperature. The solubility was determined dissolve about 95% of the solid. The boiling mixture was filtered, and water then added to the filtrate until a second after sufficient time of agitation a t 60" t o permit liquid phase appeared. Precipitation commenced quickly. establishment of equilibrium, but before any deAfter being cooled to room temperature the precipitate was composition could be observed. The estimated filtered and dried. This recrystallization procedure was precision is, of course, quite low. carried out twice. By this method, that portion of the
sample least soluble in butanol is discarded, and the portion most soluble in water is not, recovered. No trace of chloride ion could be detected in samples of the purified soaps. Throughout this procedure, protection from atmospheric COZwas necessary, especially in work with calcium oleate. Final drying waR in uacuo at 61' for several hours. The X-ray diffractionpatterns of the soap samples were assigned to the anhydrous phase. Yields after the over-all preparation and purification procedure were good with the exception of calcium oleate. Saturated Solutions.-Each saturated solution was prepared by agitation of an excess of anhydrous solid in a large Pyrex stoppered bottle about */4 full of water (co;ductivity grade). Runs were made a t 26.7 and 60.0 & 0.6 . Agitation periods of from two days to six weeks were involved. Attainment of true solubility equilibrium was indicated by constancy of calcium ion concentration on successive leaching cycles, independent of time of additional agitation and of amount of excess solid phase present, as described below. Analysis.-Samples withdrawn from the bottles a t various time intervals were filtered (fine paper). Aliquots of the filtrates were evaporated in planchets, in triplicate, for measurements of the radiation. Suitable standards were prepared from standard solutions of the calcium soaps in anhydrous butanol.
Results In general, after sufficient time of agitation to permit establishment of equilibrium, the calcium concentration in successive samples withdrawn from a given bottle became quite constant. It would be concluded a t first that the true solubility values had been obtained. However, a direct dependence of such apparent solubility values on amount of excess solid phase present was noted in replicate determinations. I n the study of such extremely insoluble compounds, very small amounts of relatively soluble impurities (e.g., due to a trace of C1zfatty acid in the original Clefatty acid sample) would cause the observed solubility value to be noticeably high. (2) W. Fahrion, Cham. Umschau., 23,34 (1916); J. SOC.Chem. Ind., 55, 932 (1916).
(3) W. D. Langley, M. G. Rosenbaum and M . M . Rosenbaum, J . B i d . Chem., 99, 271 (1932). (4) B. H. Kemp and l7. H. Fish, Virginia J. Sci., 1, 127 (1940). ( 6 ) N. P. Datta, J. Indian Chem. Sac., 16, 573 (1939). (6) Dr. E. S. Lutton kindly supplied the fatty acid samples. They had been purified by repeated crystallization of the methyl esters, followed by fractional distillation.
I
TABLE I SOLUBILITY (MOLESPER LITER)OF CALCIUM SOAPSI N WATER' Calcium stearate Calcium'palmitate Calcium laurate Calcium oleate
(2 7 & (2.8 & (4.21 j , (1.34 &
26.7' 1.7) X 1.8) X 0.17) X 0.2~) x
60' 10-6 (1.0 0 . 6 ) X 10-6 10-6 ( 1 . 5 1.3) X 10-6 10-6 (4.68 =k 0.29) X 10-6 10-6 (2 zt: I ) x 10-4
Only calcium oleate shows a very large change in solubility in the temperature range studied. It would appear that the Krafft point of this soap is at least being approached a t 60". Solubilities were also determined by weights of residue left on evaporation of large volumes (e.g., 3 liters) of saturated solutions. Although the accuracy of this method is very poor, the agreement of results with those given by the radiochemical method indicates that there is no metathetical reaction between calcium soaps and water, such as formation of "acid" or "basic" soaps in the solid phase. Analysis of such a calcium stearate residue gave 6.14% Ca, theory 6.4%. These molar solubility values are roughly one to two orders of magnitude smaller than previously reported values based on direct determinations. This emphasizes the importance of removing the impurities (presumably homologs of lower chain length) from the calcium soap samples. The solubility of calcium stearate at 26.7' reported here is in fairly good agreement with Datta's value5 at 35" based on potentiometric titration of stearic acid hydrosol with Ca(OH)*. Nature of the Solid Phase.-Marked changes in the X-ray diffraction patterns of calcium stearate, palmitate and laurate were observed on equilibration with water, indicating true hydration of the crystal. Thus, the solid phase in equilibrium with (7) The 'precision indicated (except for calcium oleate, SO0) represents the 95% confidence interval of the mean of points (4 to G points), calculated from the standard deviation. The points were chosen after a sufficient number of leaching cycles (4 to 6 ) to establish reproducibility, an additional time of agitation (22 to 383 hours) being used at each point.
C
NOTES
June, 1958
755
the saturated solution is the monohydrate8; ,8942 polymorphism reported by Vold, et al., for impure 1.000 soap samples, was not observed in this work. No change in the X-ray pattern of calcium oleate could be detected. The dihydrate has been re- 0.0000 ,1042 p~rted.~ The nature of hydration in this case inay .1560 be noli-coordinative lattice absorption. (8)IC. W. Gardiner, n4. J. Buwger and L. E. Smith, THISJOURNAL, 49, 417 (1945); IE. D. Vold, J. D. Grandine and M . J. Vold, J . Colloid Sei., 3, 339 (1948); M . J. Vold, G. S. Hattiangdi and R. D. Vold, ibid., 4, 93 (19.49). (9) F. Hoppler. Faltc'u. Seifen, 49, 700 (1942).
THE SYSTEM SODIUM CHLORATE-SODIUM CHLORIDE-WATER AT VARIOUS TEMPERATURES
.2600 .3917 .4702 .6158 ,7228 .7562 .8723 .9202 1.000
. 9696 ....
5.31 5.35
1.453 1.467
8.82 8.20 8.04 7.50 6.81 6.39 5.47 4.75 4.89 4.90 4.89 4.90
Temp. 45' 1.201 .... 1.226 0.0145 1.240 ,0232 1.267 ,0384 1.308 .0580 1.336 ,0807 1.398 ,1134 1.458 ,4510 1.462 ,9365 1.476 .9506 1.481 ,9756 1.491 ....
A A
1.42
..
..
I)
1.12. 1.15 1.09 0.98 1.07 1.01 1.22 1.20 1.52 1.61
B B B B B B B,A A A
A A
..
BYTHOMAS 5. OEY AND DONALD E. KOOPMAN Receaued January 1 8 , 1868
The available data2on the ternary system sodium chlorate-sodium chloride and water were found inadequate for use in the study of the yuateriiary system involving the three components of this system and sodium chlorite. The present work has been undertaken to re-examine this teriinry system over much wider ranges of temperatures and concentrations. TABLE I
',
s4 2 01
or
03
as
a4
a6
07
08
a9
X
Fig. 1.-The
NdIO,
system sodium chloride-sodium water a t 25'.
chlorate-
THETERNARY SYSTEMSODIUM CHLORATE, SODIUM CHLORIDE A N D
WATER
Solid phase: A, NaC103; B, Na41 2:
Compn. of s o h . 'W Sp. gr.
Compn. of wet residue X w
0.0000 ,1593 ,2142 .2696 ,3887 .4394 ,4722 .6175 ,6940 ,7478 ,8362 .9163 1.000
9.01 8.18 7.95 7.64 7.01 6.66 6.57 5.55 5.75 5.82 5.79 5.82 5.88
Temp. 25" 1.200 ,... 1.240 0.0165 1.255 ,0290 1.271 .0337 1.309 .0604 1.327 .0567 1.340 ,0548 1.402 .4689 1.408 ,9341 1.414 .9476 1.423 .9669 1.429 .9844 .... 1.440
0.0000 ,0948 ,1808 .2265 ,3333 ,4382 ,5932 .6754 .7060 ,8133 ,8659
8.96 8.48 8.03 7.79 7.22 6.62 5.67 5.14 5.18 5.26 5.29
Temp. 35" 1.201 .... 1.224 0.0130 1.246 ,0310 1.259 .0350 1.289 ,0497 1.325 ,0781 1.388 .1066 1.430 .6018 1.433 .9112 1.444 ,9466 1.451 .9599
Solid phase
.. 0.85 1.04 0.95 1.10 0.88 0.76 3.39 1.36 1.21 1.17 1.09
.. .
I
1.12 1.35 1.20 1.05 1.17 1.00 1.40 1.55 1.50 1.55
.
B B B B B B B B,A A A A A A
' {
0.1
NaCi
a2
a3
a4
n
as
a6
a7
as
as
NaCI03
Fig. 2.-The
system sodium chloride-sodium water at 35".
cldorate-
Fig. 3-The
system sodium chloride-sodium water at 45'.
chlorate-
B B B B B B
-
6.4
A A A
(1) This investigation was made under a grant from the National Science Foundation (NSF-G2750). ( 2 ) F. Winteler, 2. Elektrochem., 1, 360 (1900); J. Billiter, Monatsh., 41, 287 (1920); C. DiCapua and U. Soaletti, G a m . chim. dd.,51, 391 (1927).
The system sodium chlorate-sodium chloride and water is a simple one; no double salts have been found within the temperature interval 25-45', the solid phases found in the equilibrium mixture being sodium chlorate arid sodium chloride. Materials.-The sodium chlorate and sodium chloride used were of the analytical reagent grade. The impurities
