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March, 1962 on the rotatable turret. The lithium was introduced into the molten salt by means of a perforated stainless steel basket which was attached to a stainless steel rod and which acted both as a container for the lithium and a stirrer. The basket was removed after equilibration at temperature with stirring, and a sample of the molten salt-rich phase was taken by means of a sampling device described pre~iously.~(Sufficient additions of lithium were made to ensure the presence of two liquid phases at the equilibration temperatures.) The salts were held in both molybdenum and stainless steel crucibles, and no dependence on container material was noted in the measurements. The temperature limit of the apparatus wa,s 1000”. Attempts to determine the metal-salt phase equilibria in the Li-LiC1 and Li-LiI systems by thermal analysis above this temperature were unsuccessful. Materials.-The LiF was optical grade single crystal material (Harshaw). LiCl and LiI were reagent grade materials which were purified by very slowly heating in the presence of HCl and HI, respectively, to just below the melting temperatures of the salts, melting under dry argon, and filtering while molten. The Li was analvzed for other alkali metals, as well as C, Nz, and 02 and was found to contain less than 0.05 mole % ’ of impurity.
NOTES
573
THE LIQUID-LIQUID SOLUBILITY OF CYCLOHEXANE AND PERFLUOROTRIBUTYLAMINE AT 25’ BY RYOICHI FUJISHIRO AXD J. H. HILDEBRAND Department of Chemistry, University of California, Berkaley 4, Calzfornia Received October $4, 1961
The usual method for determining liquidliquid solubility is to observe visually the separation of a mixture of known composition into tn-o phases as the temperature is lomered. This can be done accurately only in the region around the critical point, and most composition-temperature curves have not been carried very far down the descending branches. The reliable portions of available curves for different liquid pairs extend over different ranges of temperature, and it is difficult to make a systematic comparison of such systems at a common temperature. Parnmeters calculated from critical temperatures are Results and Discussion unsatisfactory also because the structure of mixThe results for the Li-LiF system are shown in tures near the critical point is extremely complex, Fig. 1. The data obtained by the ball check valve and not amenable to model treatments that are method for the salt-rich phase are not as reliable6 reasonably applicable outside this region. It is as those obtained by the decantation method for very desirable to have figures for liquid-liquid the metal-rich phase; therefore, the salt-rich por- solubilities a t a standard temperature, preferably tion of the diagram is shown by a dotted line. a t 2 5 O , by determining the composition of both However, both methods give re,sults which fit very phases by analysis, We present the equilibrium compositions of the well with the data from thermal analysis. Figure 1 shows a comparison of all five alkali two-liquid phase of the system cyclohexane metal-fluoride systems. The regular trend in the n-perfluorotributylamine, analyzed by aid of the miscibility of the alkali metals with their molten large difference in their densities. We obtained fluorides, which is representative of the trend in the densities of the pure components, of solutions all the alkali halides, is very apparent. This trend of known composition, and of both saturated is st rapid increase in miscibility with increase in phases. Our figures are given in Table I. atomic number of’the alkali metal, that is, with the decrease in internal pressure8 of both salt and metal. TABLE I The solubility of lithium in lithium chloride was DENSITIESAT 25’ AND MOLEFRACTIONS OF CBH12, z,, AND found to increase from a value of 0.5 st 0.2 mole % (CZs)J”, 2 2 a t 640’ to 2.0 k 0.2 mole % a t 1000’. The soluUnsaturated bility of lithium in lithium iodide increased from a 21 d d value of 1.2 f 0.5 mole yo a t 550’ to 2.5 f 0.5 0 1.8714 0 0.7741 mole % a t 950’. These values are the lomest found 1.8131 0.00264 0.1399 ,7822 among the alkali metal solubilities in their chlo1.7972 .1828 rides and iodides. .2025 I ,7884 It has been demonstrated that the Li-LiF system as wc4 as the measured portions of the LiSaturated LiCl and Li-LiI systems follow the trend discussed Phase A Phase B above. Therefore, approximate delineation of the 0.218 1.7827 0.00317 miscibility gap for the Li-LiC1, Li-LiBr, and Li11.7819 LiI systerns can be deduced from comparison with the other alkali metal-halide The specific volumes coresponding to these I n light of this, further phase studies in the djfdensities plotted against mole fractions of the unficult lithium systems were not performed. saturated solutions give straight lines which, Attempts to measure the electrical conductivity extrapolated the short distances to specific volumes of the lithium systems mere unsuccessful, be- of the saturated phases, give the mole fractions in cause of reaction between the lithium solutions the two equilibrium phases, A and B. and the synthetic sapphire or single crystal magThe difference between the solubility parameters nesia capillary cells used in our conductivity ap- of the pure liquids, 61 and 82, in the approximate paratus.’ No insulating material has as yet been solubility equation found which will withstand attack by these soluIn al = In z1+ vI&(S2 - S#/RT tions. (1) (V denotes molal volume and 4 denotes volume (8) J. H. Hildebrand and R. L. Scott, “The Solubility of Nonfraction) Can be calculated by aid O f the following Electrolytes,” Third Edition, Reinhold publ. carp., N~~ York, N. y., 1960. relations
+
2)
Noms
574
(a) the four equations of the above form for aI and at in phase A and in phase B. (b) a l A = UIB, UZA = UZB, X?B = 1 (c) 516 XZA = 1, 51B VIA ‘PzA = 1, (PIB f ‘92B = 1
+ +
+
By combining these, one may obtain the equatioii
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(2)
Rotariull at the suggestion of the senior author, applying the same reasoning, but starting with ai1 equation different from eq. 1, derived a similar equation, but one in which composition is expressed in volume fractions only, now known to be inferior to eq. 1. Substituting into this equation the molal volumes VI = 108.5 cc., ~7~ = 359 cc., the mole fractions according to Table I, and the corresponding volume fractions, gives & - 6 2 = 3.08. The &values from energies of mporization give 61 8 2 = 8.2 - 5.9 = 2.3. This mixture adds another to the many examples2 of mixtures of an aliphatic hydrocarbon and a fluorocarbon that are mutually less soluble than their solubility parameters would indicate according to eq. 1. I n phase B, x1 is 0.997, and a1 = I in both phases, hence we can calculate 8 2 - 61 for phase A by means of ey. 1. The result is 3.02, practically identical with the figure aboye, calculated by eq. 2. This work has been supported by a grant from the National Science Fouiidation. (1) G. J. Rotanu, R . J. Hanrahan and R. E. Fruin, J . Am. Chem. Soc., 76, 3752 (1954). (2) (a) J. H. Hildebrand, J . Chem. Phpis., 18, 1337 (1950); (b) J. B. Hickman, J . Am. Chem. Soc., 77, 6154 (19.55).
ULTRACENTRIFUGAL DETERJIINATION OF T H E MICELLAR CHARACTER OF KOK-IONIC DETERGEST SOLUTIOKS. I11 BY C. IT. DWIGGIKS, JR.,AND R. J. AOLEN Baitlesvzlle Petroleum Research Center, Buleau o f Mznes U. S. Department of the Inteisor, Bartlesczlle, Oklahoma Recstved October 66,1961
Determiiiatioiis of micellar molecular weights of non-ionic detergents in aqueous solutions by ultraceiitrifuga1132and light-~cattering~-~ methods have shown t>hat micellar molecular weights are highly dependent upon the types of detergents studied. In addition, ultracentrifugal studies have shown that the micellar molecular weights are highly dependent on the temperature2of detergent solutions. Thus it is likely that micellar molecular weights are dependent on the ethylene oxide chain length of the detergent molecules. The usual polyoxyethylated alkylphenol detergents are available as high-purity surfactants, but such detergents usually have a rather wide (1) C. W. Dwiggins, Jr., R. J. Bolen and H. N. Dunning, J. Phye. Chem., 64, 1175 (1960). (2) C. W.D-tggins, Jr., and R. J. Bolen, sbzd., 66, 1787 (1961). (3) 4. 51. Mankowich, zbzd., 68, 1027 (1954). (4) P.Debye, zbad., 61, 18 (1947). ( 5 ) P. Debye, J . Appl. Phye., 16, 338 (1944). (6) M. J. Schick, F. R. Eirich and S.M. Atlas, ”hIlcellar Structure of Nonionio Detergents,” 139th Kational Meeting, Amerlcan Chemical Society, St. Louis, Missouri, March, 1961.
Vol. 66
distribution of numbers of ethylene oxide groups. Polyoxyethylated alkylphenol detergents having no chain length distribution are very difficult to obtain, but fractions having much narrower chain length distributions than the detergents available commercially may be obtained by molecular distillation.’ Several properties of molecular distillation fractions of a polyoxyethylated nonylphenol detergent were ~ t u d i e d ,and ~ physical properties were significantly different for each detergent fraction. After these fractioiis became available, it was decided to study the effect of ethylene oxide chain length distribution on micellar molecular weights of polyoxyethylated nonylphenol nonionic detergent. Experimental The transient-state methods-10 for molecular weight determination using the analytical ultracentrifuge aided by synthetic boundary experiments,lI and application of the theory to determination of niicellar molecular weights of non-ionic detergents in aqueous solutions were discussed in the first tx-o papers of this series and will not be repeated. One set of detergrnt fractions was obtained from a polyoxyethylated nonylphenol having an average ethylene oxide mole ratio of 9.5. These fractions were cycles 9, 10 and 11 as described by Mayhew and Hyatt7 and had average mole ratios of 8.9, 9.7 and 10.7, respectively. A second set of fractions was obtained from a polyoxyethglated nonylphenol having an average ethylene oxide mole ratio of 6.0. These fractions were cycles 12 and 13 as described in ref. 7 and had average mole ratios of 7 . 6 and 7.7, respectively. The detergent fractions have much narrower polyoxyethylene chain length distributions than the parent detergents, but it is not claimed that they have no chain length distribution. The ultracentrifuge experiments were performed in a 2.5’, 12-mm. double sector cell, and a matched capillary synthetic boundary cell as described.’,? Great care was taken to maintain constant temperature and precise alignment, to prevent convection, and to prevent biological contamination. Values of the corrected concentration gradients at the airliquid meniscus vere used to determine micellar molecular weights.1 Pycnometers were used to obtain partial specific volumes as described previously.112 The pycno:eter water-bath was maintained constant to within 0.005 . Least-squares calculations were used to evaluate partial specific volumes as described previously.1.2 The average deviation of individual specific volume data from the lines fitted by least squares ranged from 0.00001 to 0.00004.
Results and Conclusions Micellar molecular weights and partial specific volumes for the various fractions are listed in Table 1. Decreasing the number of ethylene oxide groups results in increased micellar molecular weights in aqueous solutions. It is of interest that, for the lower ethylene oxide mole ratios, very small changes in the ethylene oxide mole ratio can produce quite large changes in micellar molecular weights. Cloud-point temperatures are reduced greatly when ethylene oxide mole ratios are decrea~ed.~ It appears that cloud-points are dependent on the micellar character of the detergent solutions, but the correlation between cloud-point temperatures and micellar molecular weights may be the result of other, more basic factors. ti’) R. L. illaybew and R. C . Hyatt, J . Am. Od Chembsts Soc., 29, 357(1952). (8) T. J. Archibald, J . Phys. Chem , 61, 1204 (1947). (9) H. K. Schachman, “Ultracentrifugation in Biochemlstry,” Academio Press, New York, N. Y..1959, PP. 181-199. (10) J. 31. Peterson and R. M. hIazo, J . Phys. Chem., 66, 566 (1961). (11) S. 111. Klainer and G. Kegeles, %bid., 69, 952 (1955).
