572
NOTES
Figure 2 shows the effect of pH on the electrical resistance of the polyampholyte membrane. Curve 1, measured in 0.1 I buffered sodium chloride, exhibits an order of magnitude rise in resistance as a function of pH. The Ionics membranes were similarily measured and practically no effect on the resistance was observed. The resistance of the polyampholyte membrane is low a t pH values at the extremes of the range used, since at these pH values the concentration of fixed charges (positive at low pH and negative a t high pH) with their associated gegenions is a t the maximum. As the pH of 6.1 is approached from either side, the concentration of charges decreases and the resistance rises. At the isoelectric point, fixed (immobile) zwitterions exist and free charge carrying gegenions are at a minimum. Another factor responsible for the trend illustrated in Fig. 2 is the choice of solute. Curve 2 shows the data obtained with 2% NaPS-0.002 I buffer solutions. Except for the first point, the resistance falls with increasing pH and does not show the sharp rise in resistance a t or near the isoelectric point. This behavior can be explained on the basis of steric considerations. At a pH of about 4.7-4.8, the polyampholyte membrane bears fixed positive charges which hinder the transport of cations across it. Consequently, current is carried by anions. However, a large fraction of the free anions is hindered and the resistance is high. As the pH is raised, the membrane charge diminishes and the Donnan exclusion effect becomes less pronounced. The transport number of sodium ions across the membrane becomes greater and the resistance falls. At the presumed isoelectric point there is no barrier to sodium ions in contradistinction to that at lower pK. As the p H is increased further, the membrane becomes negatively charged and the transport number of the sodium ion approaches unity; thus, a larger current can be maintained and the resistance continues to fall. It is conjectured that the initial rise in resistance is due t o a decrease in the gegenion concentration (and fixed ion concentration) but not enough of a decrease to abate the Donnan effect-that is, sodium leakage remains unchanged and minimal. This is borne out by reference to the concentration potentials a t the various pH values. As can be seen from Fig. 1, a t pH 4.32 and 4.79, the concentration potentials are close to the maximum. The slope becomes substantial a t pW values higher than 5.0. Up to this pH, ion leakage is small and the resistance is governed by the gegenion concentration. Above pH 5.0, leakage is enhanced and a greater fraction of ions becomes available for carrying current. MISCIBILITY OF METALS WITH SALTS. VI. LITI-IIUM-LITHIUM HALIDE SYSTEMS' BYA. S. DWORHXN, H. R. BRONSTEIN, AND M. A. BREDIQ Chemistry Division, Oak Ridge National Laboratory, Oak Ridge, Tennessee Received October $0,1961
The miscibility of the alkali metals with their halides was shown to decrease rapidly in going
Vol. 66
from the cesium systems to the sodium This report covers the lithium systems which were expected to exhibit a continuation of the trend; that is, low solubility of metal in salt and high (consolute) temperatures for complete miscibility between metal and salt. Experimental Apparatus and Procedure.-The points on the Li-LiF phase diagram (Fig. 1) designated by open circles were obtained by means of thermal analysis (cooling curves) as described previously .* A platinum wound Marshall furnace,
4300
Li,F - Li 4200
14oc
1ooc "C
. 9oc
80(
70(
1
CSF - CS 60(
I 25 50
\
MOLE % METAL,
Fig. 1.-Liquid
metal-salt phase equilibria in the alkali metal-fluoride systems.
stainless steel sample capsules, and a Pt-PtRh thermocouple were used at temperatures up to 1400'. At lower temperatures and at low salt or metal concentrations, where the thermal halts for liquid-liquid separation become unobservable because of the small temperature dependence of the solubility, separate sampling of the liquid metal-rich and salt-rich phases after equilibration2 yielded the points designated X (Pig. 1). For the Li-LiC1 and Li-LiI systems, an apparatus was built which permitted the direct sampling of the solution at temperature, as in some measurements on the sodiumsodium halide systems.6 The apparatus was a smaller, modified version of the conductivity apparatus used for the alkali metal-halide systems,? and has only one entry port (1) Work performed for the U. S. Atomic Energy Commisaion at the Oak Ridge National Laboratory, operated by the Union Carbide Corporation, Oak Ridge, Tennessee. (2) M. A. Bredig, J. W. Johnson, and W. T. Smith, Jr., J . Am. Chem. Soc., 11,307 (1955). (3) M. A. Bredig, H. R. Bronstein, and W. T.Smith, Jr., ibid., 77, 1454 (1955). (4) J. W.Johnaon and M. A. Bredig, J . Phys. Chsm., 62,604(1958). (5) M.A. Bredig 5nd H. R. Bronstein, ibid., 64,64 (1960). (6) M.A. Bredig and J. W. Johnson, ibid., 64,1899 (1960). (7) H.R. Bronstein and M. A. Bredig, J . Am. Chsm. SOC.,80, 2077 (1958).
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 ’ of impurity. than 0.05 mole %
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)
