2870
P A U LD. SCHETTLER, JR.,A N D ANDREW PATTEKSON, ,JIL
Liquid-Liquid Phase Separation in Alkali Metal-Ammonia Solutions. 11.
Sodium with Added Sodium Iodide
by Paul D. Schettler, Jr.,I and Andrew Patterson, Jr. Contribution .Vo. 1741 f r o m the Sterling Chemistry Laboratory. Yale University, S e w Haven, Connscticut (Received August 6 , 1.968)
Results are presented for the effect on the phasc separation of solutions of sodium in liquid ammonia of adding varying amounts of sodium iodide. The measurements were perApparatus is described which makes possible formed a t -33.35, -56.50, and -75.00'. the convenient preparation and separation of these solutions under conditions guaranteeing their long life and freedom from undesired inipurities, plus the analysis of both phases for all components. The results are discussed.
Introduction Studies by Cubicciotti2*and SienkoZbhave indicated a marked effect on the temperature of phase separation of the addition of sniall aniounts of salt in the systems Sa-SH3-NaI and K-SH,-KI. The work of these authors showed that thc addition of the iodide decreased the metal concentration in the dilute phase of the twophase solutions, and that the preponderance of the salt appeared in the dilute phase. The tendency of metalammonia solutions to decompose under any but the most extreme conditions of chemical purity and rigorous exclusion of catalytically active materials, plus the large effect of the added salt, and in all probability of any products of decomposition, prompted us to undertake a study of the phase separation in sodium-liquid ammonia solutions with added sodium iodide which is reported in this paper, with the intention of improving the manipulative details to permit analysis of both phascs for all components and to permit performance of all operations under high vacuum with precise temperature control. I n addition to accomplishing these goals, with an accompanying increase in precision of the data, we have covered a wider range of concentrations of thc coniporients and of temperature in the study than has heretofore been reported.
distillations and separations can be conducted in an inert atmosphere or under high vacuum; it is semiautomatic in its function, and effects a complete and clear-cut separation of the two phases; the solutions can be removed for analysis of all components. The operation of the device can best be discussed after a brief description of its preparation. Tube D was made from 0.25-in. medium-wall Pyrex tubing. The section a t K was made in two stages so that the very end was extremely thin. The tiny glass ball a t the end is important for two reasons: it makes accidental breaking of the tip easily detectable, and it appears to make the tip easier to break at will. At the opposite end, a constriction as small as practicable was left at L and the tube sealed to a vacuum system inside an oven. After thorough bake out, it was filled with purified helium to about 250 torr and sealed off. Glass-enclosed magnet F was immediately attached. The remainder of the apparatus was then assembled. All glass used was cleaned by soaking in potassium hydroxide in ethanol, rinsing in water, soaking in a mixture of concentrated nitric and sulfuric acids, rinsing in distilled water, and drying in an oven. The object shown in bulb C is a magnetic stirrer. When completed, the device was sealed to a high vacuum system
Experimental In 14g. 1 is shown a diagram of the glass apparatus used, incorporating several desirable features. It can
(1) This paper is taken in part from a dissertation submitted to the Graduate School, Yale University, in partial fulfillment of the requirements of the 1'h.D. degree, September, 1963.
be evacuated and baked out a t high temperature; all
M .J
T h e .lournnl of Physical Chemistry
(2) (a) D. D. Cubicciotti, J . P h w . Chem., 53, 1302 (1949); (b) Sienko, J . Am. Chem. Soc., 71, 2707 (1949).
LIQVID-LIQUID ~’HASE SEP.%HATION I N ALKALIhIETAL-AMMONIA
c
A
4
Figure 1. Schematic diagram of glass apparatus used for preparation of three-component solutions of sodium, sodium iodide, and liquid ammonia. The same device has been used for lithium-ammonia solutions. See ref. 3.
a t A inside a controlled oven, but with the distillation tube for the alkali metal, B, protruding outside into the room. The vacuum system had arrangements for triple distillation, drying, and measuring of the animonia. A weighed sample of sodium iodide (Mallinckrodt A.R.) was placed directly in bulb C, when required, before the apparatus was sealed. The alkali metal contained in a measured portion of glass capillary was sealed into tube B. The assembly inside the oven was heated to 300” under high vacuum for a number of hours to remove water from the salt and from the glass walls; a t the same time the sodium was melted out of the capillary. The apparatus was cooled after bake out and the sodium distilled from one section of tube B to the next, and finally into C. After a measured quantity of ammonia had been distilled into bulb C, the apparatus was filled with helium to about 650 torr pressure. The helium was purified with activated charcoal a t liquid nitrogen temperature. The device was then sealed from the vacuum system and placed in a low temperature thermostat. (The same device has been used for preparation of lithium solut i o n ~for ~ a study of phase separation in lithiumammonia solutions.) The thermostat permitted excellent temperature control and measurement and will accordingly be of interest to other investigators working with moderate-
2871
SOLUTIONS
sized equipment a t low temperatures. It consisted of a stirred bath of jet aircraft fuel ,JP-4 mounted inside a Revco SZC-657 ultra low temperature cabinet. A copper-lined plywood box was used for the bath. The oil was agitated by stirrers mounted externally on an insulated top made to fit the freezer. Jet fuel JP-4 remains fluid at the lowest temperatures attainable with the freezer; the viscosity increases from 0.82 cp. a t 25” to 1.1 cp. a t -75”) or from about that of water to about 40% more than that of water. The low viscosity made stirring effective a t all temperatures. The electrical insulating properties are good and are maintained in spite of the inevitable condensation of small amounts of water in the bath liquid. Temperature control was obtained by setting the deep freeze a t a value suitably below that desired in the bath and supplying intermittent heat to the bath with an electrical heater operated from a JIueller bridge in a much simplified version of the arrangement suggested by FUOSS, et aL4 The circuit diagram is given in Fig. 2 GNVALYIYOIR MIRROR lM?tE
RfO
FlSHfR
I
mriEn -mRuwwin P C 111 (POI
Figure 2. Wiring diagram for thermostat regulator. The photoconductive cells are mounted on an aluminum and plastic ref. 4) which can be hung on the block (see FUOSS, galvanometer scale to intercept the light spot but leave the spot free for visual observation of bridge balance; the block is moved to the desired position after Mueller bridge balance is obtained. When the contacts of the latching relay are in the position shown, the heater is on, and only the photocell on the left will operate the circuit. The Fisher relay must be switched to the “normally open” and “delay off” positions. The heater or heater control relay may be plugged into the “normally off” plug marked “heater in” or wired to additional contacts of the latching relay. A variable transformer used to control the heater input power is not shown here. (3) P. D. Schettler, Jr., and A. Patterson, Jr., J . P h y s . Chem., 68, 2865 (1964). (4) J. J. Zwolenik, J. Lind, and R. M. Fuoss, Rev. Sei. Znstr., 30, 575 (1959).
Volume 68,Number 10
October, 1964
2872
and the operation is explained in the legend. Temperature stability quite equivalent to that commonly possible near room temperature could be maintained, in the order of = t O . O l " or better, at any of the operating temperatures chosen. The platinum resistance thermometer and bridge were used for precise teniperature measurement. As a word of caution, JP-4 is appreciably volatile at room temperature and is prone to build up electrostatic potential when agitated, so one must guard against the fire and explosion hazard associated with its use. If the amounts of metal and ammonia used and the proporti'on of the apparatus were correct, following thorough mixing of the solution and attainment of temperature equilibrium, it was then possible to place a magnet at G to hold tube D in place, turn the apparatus upside down in the direction of the arrows, and allow the solution to run down and fill the main tube to H. Tube D was then lowered by the outside magnet, causing a sample of the less-dense phase to be displaced and run down the spiral side arm into tube I. Tip E was then sharply rapped against A with the magnet, breaking E. The pressure differential existing in the device caused tube D to fill half full of the denser phase. The apparatus was then rotated into its upright position in the direction opposite to the arrows, causing the concentrated phase to run into tube J and D to return to its original position. The remaining unseparated solution was discarded into C. Sealed tubes of each phase were then obtained by placing the apparatus in a flask of Dry Ice-alcohol and sealing off J at the constriction. The main tube of the device was then cut just above H, and tube D was removed and sealed at K. The solution was briefly exposed to possible entry of air a t this point. While reaction with the air is not of any analytical consequence, loss of ammonia was very small, since tip E was extremely fine, the net motion of helium a t the moment of opening was into tube D, and the vapor pressure of ammonia at -78" is low. The sealed tubes were allowed to reach room temperature, washed in distilled water, dried, and kept for a period until they were weighed. They were then frozen in liquid nitrogen, broken open near the upper end, and the ammonia was allowed to evaporate into a stream of purified nitrogen. The metal was then oxidized in a stream of water vapor a t reduced pressure; the pieces of glass were removed and carefully washed, dried, and weighed, and a titration conducted on the solution remaining with standard nitric acid. Titration was then made for iodide with standard silver nitrate solution. All titrations were conducted under an atmosphere of butane. The Journal of Physical Chemistry
PAULD. SCHETTLER, JR.,AXD ANDREWPATTERSON, JR.
Results The results are given in Table I and are plotted in Fig. 3. The temperatures used mere -33.36, -56.50, and -75.00". By extrapolating data at -75" to zero sodium iodide concentration, values for the phase separation of sodium a t this temperature have been obtained and are reported in ref. 3. At -56" pairs of points with no sodium iodide present were included in the determinations, also reported in ref. 3. Data of CubicciottiZaand SienkoZbare reproduced in Fig. 4, plotted in the same manner as Fig. 3 , for comparison. .DATA AT -33,35'C +DATA AT -56.50'C oDAIA AT-75.00'C
CONCENTRATED
.05
.oa
.06
a9
,IO
.I2
DILUTE PHASE
,010
,005
0
MOLE FRACTION SODIUM METAL
Figure 3. Plots of the data of Table I as mole fraction of sodium iodide us. mole fraction of sodium metal for concentrated and dilute phases. Notice that the abscissa on the concentrated phase diagram does not start a t zero mole fraction of sodium metal.
005
Lo 0 SIERLO
I-
CCNCE N I R A I E D QHASE
A i -11 I'C A I -3I'C
X CUBlCClOTTl
I 005
I 010 UOLl l R A C l l 0 l l
I 015 ALKALI 31111
I 020
Figure 4. Plots of the data of Cubicciotti28 and SienkoZb as mole fraction of sodium iodide us. sodium metal (or potassium iodide and potassium metal) for the concentrated and dilute phases. Pl'otice that the abscissa on the concentrated phase diagram does not start a t zero mole fraction of metal.
I
LIQCID-LIQVID PHASE
SEPARATIOX IS
ALKALI>IETAL-AhlMOKIA4
Table I : Phase Separa1,ion Data on Sodium-Sodium Iodide-Liquid Ammonia Solutions -----IIilute Nza
phase--------N 3
---
Concentrated phase--.V*' -3'8'
0,007974 0,007406 0.003723 0.002955 0.001975
Temp. -75.00O 0.00255 0.09890 0,006283 0,10265 0.01906 0.10656 0 02500 0.10777 0.03021 0.11072
0.0001670 0.0 0,000118 0.0000911 0,0001668
0.01706 0.10670 0.01349 0.0084O6 0,00395 0.00362 0.0006123
Temp. -56.60" 0.07589 0.0 0.0 0,075'32 0.000492 0,086421 0.01243 0.09106 0.04264 0,09937 0.10198 0.04702 0.1045 0.05240
0.0 0.0 0.0001437 0.0001886 0,0002633 0,000652 0.0010806
0.02061 0.008996 0.003552 0,001707 0,001497 0,001183
Temp. -33.35' 0.007593 0.05493 0.04664 0.07411 0.04855 0.08946 0.05833 0.09618 0.06529 0.1014 0,08767 0.1161
0.00186 0.00123 0,000793 0,001306 0.000559 0.000305
a ]V2 = mole fraction sodium metal; dVa = mole fraction sodium iodide. Unprimed quantities refer to dilute phase, primed to concentrated phase.
Discussion The curves for thc dilute phase are siinilar to each other, concave upward with negative slopes. They intercept the Ns = 0 'axis (zero sodium iodide) depending on whether the temperature is below or above the consolute temperature of the two-component diagram. At each temperature, as the sodium iodide concentration increases, the sodium concentration decreases. The concentrated-phase curves do not show such rcgularity. At -75" only a very small amount of silver nitrate was required to reach the iodide end point, the results are scattered, the experimental uncertainty in the iodide concentration is large, and no consistent trend of the results is apparent; the sodium concentration remains constant within lo%, although the concentration is comparatively large and tends toward larger values as thc sodium iodide concentration in the dilute phase increases. At -56" the amount of sodium iodide in the concentrated phase shows a definite increase with increasing total sodium iodide and the sodium concentration increases as well. At -33" this trend is partly reversed, for the sodium iodjde concentration decreases as the total added iodide concentration increases, but, there is a relatively large increase in the sodium concentration.
2873
SOLUTIOXS
An attempt at interpreting these data should start, ideally, with a calculation of the concentrations of all t'he species present' in each phase and an analysis of t'he equality of the chemical potentials of each of these in each phase as t'he criterion for the possible existence of two phases in equilibrium wit'h each other. The number of species present in each phase is considcrable, including, according to a recent proposed model,5 the following: KaO, S a + , S a z o , e-, and (Nae-)-. All these species are in equilibrium with each other via three equilibrium reactions. Equilibrium constants for the interactions between these several species have been derived by Arnold over a range of temperature, although the precision and the limited temperature range of the experimental data used in calculating them restrict the number of significant figures in these constants and t'he temperature range over which they may be meaningfully applied. I n principle, however, it is feasible to calculate the concentrations of all species in each phase, with due respect to the activities of t'he species. Practically, this can be done only for the dilute phase, and there with limited precision, for the concentrations are nearly ten times larger than those for which the calculation may be expected to be valid. The concentrated phase is far too concentrated to permit a meaningful calculation, although it' is quite clear that in both phases t'he concentrations of the M + ion, XIz dimer, and diamagnetic 11' center are decidedly larger than those of the electron arid the neutral atom. Sodium iodide is appreciably ionized in ammonia solution, though not completely so, and data on its degree of ionization a t these concentrations are unavailable. Having all these data a t hand, however, it still seems clear that there are at least a large, if not infinite, number of values of these equilibrium concent'rations which might satisfy the requirements, whatever they may be, for the simultaneous existence of two liquid phases. The only experimcntally available quantities are Saotota1 and I- in cach phase. I n the face of these several difficulties, some approximations can be attempted which may guide further st'udy. As a first approximation, Onsag& has suggested that a solubility product expression might correlate the data. While it is apparent from the discussion of the first paragraph that this will not hold at all temperatures studied, a t -75" thc expression [Na+][e-] = (5) E. Arnold and A. Patterson, J . Chem. PhUs., accepted for publication; E. Arnold, Dissertation, Yale University, 1963. Arnold's terminology for the centers present is M, M +, M2,e, and M ', respectively. (6) L. Onsager, private communication.
Volume 68, Number 10
October, 1964
2874
k, or the equivalent ( N z
PAUL
+
D.
SCHETTLER,
JR.,AND ANDREWPATTERSON, JR.
N 3 ) ( N 2 can ) be calculated, continuous trend from smaller to larger values of this assuming complete ionization and lack of any associaproduct, of which note should be taken below in refertive reactions. Refer to Table 11, in which are comence to column E. pared a number of products and quotients of ion conThe general dissatisfaction which one feels with this centrations; column A contains data calculated as just approach need hardly be remarked. Although it is suggested. I n such a n approach the concentrated true that activities and possible association reactions phase is regarded as analogous to the solid phase of a have been neglected, it is possibly more serious to slightly soluble salt, and the ion product is written for neglect the suggestion that the value of [Xa+][e-] the species in solution. As common ion is added, should be equal in the two phases, rather than equal to a constant, as has just been assumed. That is, if sodium "precipitates" into the concentrated phase. the activity coefficients do not change too rapidly with The concentration products so obtained vary from the concentration, and barring complications from the mean by +22 to -24%. If the second and fifth points many solution species present, the activities should be are omitted as excursions beyond experimental error equal in the two phases and the ratio of the concentra(they cancel each other out in the calculation of the tion products [Na'] [e-] in the two phases should equal mean), then the variation is + 2 to -0.6% from the a constant. Column C, representing the quotient of mean. I n the concentrated solution in equilibrium with the dilute phase, column B, the influence of XU'~ columns B/A, is a test of this possibility. The quotient is small even though the data are scattered, since K3 varies more severely than does either colunin A or B is much smaller than N2. As shown in B, there is a if the five points in column A are included, or about the same as column A if the second and fifth points are excluded. If one takes the theoretical calculations of Arnold, Table I1 : A Variety of Ion Product and Quotient which have been performed for -33.5' as the lowest Calculations for Phase Separation temperature owing to the scarcity of experimental A. Dilute solution B. Concentrated C. (Ion prod.),,,,d/ data as noted above, extrapolates his graphs to the solution (ion pTod.)dil higher concentrations represented in the solutions here (N2 + N 3 ) ( N d a (NZ + N 3 ) ( N d 5 (B/A)" dealt with, and reads off values of [Na+] and [e--] 8.38 x 10-5 9 79 x 10-8 1.17 X 10* 10.1 10.53 104 corresponding to the concentrations of total sodium 8.47 11.37 1.34 found in this paper (see Table I) at -75", and calcu8.27 11.62 1.40 6.33 12.27 1.93 lates their products, one obtains the values shown as Mean: 8.32 X Mean: 11.12 X 10-8 Mean: 1.38 X 10% columns D1 (dilute phase) and E (concentrated phase). Dev.: +22, -24y0 Dev.: f 1 0 , -12% Dev.: 1-39, -24% The products found in column Dz are obtained by Dl, D z . ~Dilute solution E. Concentrated FI, Fa. (Ion prOd.),o,,d/ adding to [Na+] the additional concentrations of Na+ solution (ion pmd.)dil [Na+l[e-ItheorC [Na. +I [e-ltheorc (E/DI, DdtheorC found in solution if the sodium iodide were wholly ionized in the solvent and multiplying by the same 7.56 X 10-2 3.43 1.48 X 10' 5 09 x 10-8 2.20 7.94 4.07 0.91 8.74 1.95 [e-] as in column D1. The quotients of columns 8.32 10.9 0.50 0.761 1 6 . 8 D, and D, divided into the theoretical ion products 14.7 0.46 8.51 0.578 18.6 8.77 23.7 0.48 0.370 1 8 . 1 for the concentrated solutions (E) are given as columns Mean: 1.17; 13.4 X 10-3 Mean: 8.22 X 10-2 Mean: 11.4; 0.77 X 10 F1 and Fz. One may justify the use of data a t Dev.: +88, -68%; Dev.: +6.7, -8.0% Dev.: +108, -70%; + 92. -40% +62, -24% -33.5' by noting that there are no inversions or changes of slope of the equilibrium constants as a function G. (Ion prod.)thoorc/(ion prod.)oonodd of decreasing temperature (see ref. 6) ; accordingly, (ED) 7.72 although the concentrations are in error, their trend 7.54 as a function of concentration is not markedly incorrect, 7.32 7.32 and since they are used in ratios as normalizing func7.12 tions only, no harm is done if the fact is kept in mind Mean: 7.41 that the theoretical and experimental data are for difDev.: f4.2, -3.5% * Column Dz ferent temperatures. a These data are derived from Table I a t -75". Colunin Di thus obtained varies widely; it must be includes [Na+] from the added sodium iodide, assuming cornkept in mind that in the solution under consideration, nlete ionization. Theoretical values are all for a temperature of -33.5", and are taken from Arnold, see ref. 5. The for which NaOtotalhas been determined experimentally, theoretical Ion product is for calculations at -33.5" while the decreasing because sa+ mas being this quantity experimental data are for -75". added in the form of sodium iodide. To compensate The Journal of Physical Chemistry
LIQVID-LIQVIDPHASE SEPARATION IN ALKALIMETAL-AMMONIA SOLUTIONS
for this, column Dz has been calculated. The variation in the product is here somewhat less, but the trend is toward larger values of the product with increasing NB. I n the concentrated solution, column E, the behavior of which should correspond closely to that of column B, the trend is toward larger values of the product as the sodium iodide concentration in the dilute phase increases. Column F, is a numerical test of the way in which colurnn C might have been expected to vary if our rather inadequatc theoretical guidance were significant. The direction of the variation is reversed and the values of Fz vary much more than do those in C. Were activity data available for the sodium iodide, this situation would be improved. Finally, in column G are shown the quotients of the data in column E, theoretically derived and without complications of sodium iodide being present, divided by the data in column B, which are experimental and essentially (because of the very small amount of sodium iodide in the concentrated phase) under the same circumstances. These quotients are substantially constant, appreciably more so than any of the raw data found elsewhere in Table 11. This is not surprising, since this is the situation in which the assumption of constancy of activities and noninterference in the complex equilibria by the added sodium and iodide may be expected to hold. Similar calculations have been performed for the other temperatures, but the data are even less consistent, and, of course, there is no theoretical guidance tlo explain the inversion of the order of the values of Na' a t -%and -33". Franklin and Kraus' have reported measurements of the degree of ionization of sodium iodide a t -33" derived from conductance studies. The range of con-
2875
centrations covered does not extend as high as those involved in column D1,nor are the data given for - 75" ; however, to demonstrate that if activity data were available considerable improvement would result in the character of calculations such as those in Table 11, we have recalculated column D2 with activities extrapolated from Kraus's data, ranging from degrees of dissociation of 0.36 to 0.15. Instead of values of Dz between 5.09 and 18.1 X mean 13.4 X with deviations from the mean of $62 to -24%, we obtain values between 3.85 and 3.03 X mean 3.41 x with deviations of +13 and -11%. Not only js the constancy of quotient columns E and B of interest, but it should be observed that the trend in each is toward larger values of the product [Na+]. [e-] as the concentration of sodium in the concentrated phase rises. Constancy of this product is not to be expected in either dilute or concentrated phases, owing to the interplay of the equilibria between other solute species also present. Indeed, it does not appear from Arnold's results that any other combination of concentration product functions will remain constant, for [e-] is increasing in this range of concentration more slowly than any other quantity. Accordingly, until one can compute or determine activities, any approach such as attempted in Table I1 is bound to fail. Further, in a system with as many degrees of freedom as here represented, it is not likely that each activity in one phase must be identical with that of the same specific solution species in the other phase.
Acknowledgment. This work was supported by the Yational Science Foundation. (7) E. C. Franklin and C A. Kraus, *J. Am. Chem. SOC.,27, 191
(1905).
Volume 68,Number 10
October, 1964
