3535
LIQUID-LIQUID PHASE SEPARATION IN ALKALI RfETAL-AMMONIA SOLUTIONS
of the Texas -4& -11 University System have been used extensively in the course of this research. Dr. Roger
D. Whealy has kindly supplied us with samples of the compounds studied.
Liquid-Liquid Phase Separation in Alkali Metal-Ammonia Solutions. 111. Sodium with Added Sodium Bromide and Azide
by Patricia White Doumaux' and Andrew Patterson, Jr. Sterling Chemistry Laboratory, Yale Unioersitv, New Haven, Connecticut Accepted and Transmitted by T h e Faraday Society
06620
(February 20, 1.967)
Experimental data are given for the effect on phase separation in sodium-liquid ammonia solutions of adding varying ?mounts of sodium salts. Xeasurements were made a t -32.90' with sodium azide and at -32.90 and -56.39' with added sodium bromide. The data are compared with previously reported measurements on sodium iodide. The additions all raise the consolute temperature, broaden the temperature-composition curve, and alter the distribution of components: at a given temperature as the total salt is increased, the salt migrates to the dilute phase and the sodium metal to the concentrated phase. The magnitude of these effects varies, approximately linearly, with the amount of salt added, but depends also on the salt used in the order KaI > N a y 3 > SaBr, the most effective listed first.
Introduction In a previous paper,* the effect of adding sodium iodide on liquid-liquid phase separation in sodiuniammonia solutions was examined. It was not found possible to explain the results in terms of a simple common ion effect. On extending these measurements to include salts with other anionic constituents as reported in this paper, one finds that although the concentration of the salt is a significant factor, it is principally the anion of the salt which influences the results. This finding is further supported by data on phase separation in the sodium-potassium-ammonia system which are reported separately.3
first two reacted with the metal-ammonia solution, and the remainder were too insoluble to bring about phase separation a t temperatures of -33" and below for which data on sodium iodide were available for comparison. Sodium bromide and azide suffered none of these failings; they were studied in detail by a procedure which was essentially that of Schettler and Patterson,2 with minor modifications. Samples were taken of the separated phases, the ammonia was quantitatively removed, and aqueous solutions of the residues were analyzed for base and salt by successive titrations (1) This paper contains material taken from a dissertation submitted
Experimental Section
by P. W. Doumaux to the Graduate School, Tale University, in partial fulfillment of the requirements of the degree of Doctor of
In preliminary measurements, sodium thiocyanate, sodium sulfide, sodium chloride, sodium cyanide, and sodium tetraphenylborate additions were tested. The
Philosophy, Sept 1966. (2) P. D. Schettler and A. Patterson, J . Phys. Chem., 6 8 , 2870 (1964). (3) P. W. Doumaux and A. Patterson, ibid., 7 1 , 3540 (1967).
Volume 7 1 , S u m b e r 11 October 1967
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PATRICIA WHITEDOUMAUX AND ANDREWPATTERSON, JR.
with nitric acid and silver nitrate. Because of the dissimilar properties of the silver precipitates, different equivalent point determinations had to be used for each. The precision of the analyses has been tested not only by examining the reproducibility of each of the silver determinations, but also by studying the over-all precision possible as a consequence of all the sample and solution manipulations. A series of samples containing metal, sodium chloride, and ammonia mere analyzed. The chloride-containing solutions do not separate into two phases, so the partitioned samples obtained with the Schettler apparatus were necessarily identical. The deviations in the analyses from all manipulative steps were found to lie within 1-2%, usually lye. As a further check on the azide analyses, a series of titrations were compared with gravimetric determinations of silver azide, the results falling well within these limits. Details of the handling of the different salts were as follows. Sodium bromide was dried for 5 hr in an Aberhalden pistol using toluene, stored over Drierite, and weighed in air. Both phenolphthalein and differential platinum electrode end-point determinations were made. Bromide was determined with silver nitrate using the same electrodes. In only one determination, that listed last at -56.39' in Table I, 0.01 in silver nitrate was used instead of the 0.1 in solutions otherwise employed. Reagent grade sodium azide was dried
Table I : Liquid-Liquid Phase Separation Data on Sodium-Ammonia Solutions with Added Sodium Saltb ------Diiuti~
xlr,aa 0 0 0 0 0 0 0 0 0
phase----
------Coned
A-hRX
a
phase-----
Xvax
Temperature, - 32 90" ; Salt, Sodium Bromide 0.0660 0.00215 0 0197 0.0656 0.00224 0 0208 0,00227 0 0209 0 0633 0.00239 0 01.51 0 0612 0 0120 0 0588 0 00241 0 0307 0 00295 0 00747 0 0472 0 00292 0 0031.i 0 00544 0 0399 0 00404 0 00392 0 0330 0 00333
0128 0131 0134 0176 0196 02.53 0286 0309 03.5.5
Temperatiue, - .i6.39", 0 0134 0 00311 0 0163 0 00156
0 0 0 0
A'Na
Temperature, -32 90"; 00737 0 0340 0114 0 0244 0211 0 00856 0228 0 00752
Salt, Sodium Bromide 0 0i63 0 000283 0 0764 0 0000807 Salt, Sodium 0 0734 0 0686 0 0538 0 0524
All concenl rations are mole fractions.
T h e Journal of Physical Chemistry
Azide 0 00148 0 000986 0 0019.5 0 00230
a t 50' for 21 hr and weighed in air. The glassware baking-out procedure normally used was modified to limit the temperature to 50', but this was followed by prolonged evacuation under "stick-vacuum" conditions before the solution was prepared. It was found that azide samples invariably exploded on being warmed t o room temperature for weighing, but that an equally satisfactory analysis for ammonia could be made by absorbing it in standard 1 in nitric acid in a gas washing bottle and titrating the excess with 1 in sodium hydroxide. A silver wire differential electrode had to be used for the azide analyses, both an adsorption indicator and the platinum electrodes being unsatisfactory. The chloride analyses were made with dichlorofluorescein indicator. The other manipulations mere made in the same way as described earlier;* the temperatures were the same. It should be noted that the temperature -32.90' is in fact the same and should be substituted for that reported by Schettler and Patterson2 as -33.35'. The -32.90' figure is correct, the lower figure having been in error owing to the use of a reversed sign in correcting Jlueller bridge resistance readings.
Results The bromide and azide determinations are listed in Table I. Figures 1 and 2 display mole fraction in the dilute phase us. mole fraction in the concentrated phase for sodium (Figure 1) and salt (Figure 2) as a function of temperature and salt used; the data for iodide of ref 2 are included. Figure 3 is a plot of mole fraction salt L'S. mole fraction sodium for both phases, the type of plot used in ref 2 , lvhere the plot was split into two parts of different scale. The data are crowded together on a plot made to the scale used here, but on a suitably expanded scale the tie lines joining corresponding points can be extrapolated to a consolute point for metal and salt. The results, at -32.90', are N x a = 0.0369, NSaI = 0.00235; N X , = 0.0358, = 0.00340; N x a = 0.0361, N A - ~=B0.00408 ~ (arranged in order of increasing mole fraction of added salt).
Discussion At -32.90', the results for iodide, azide, and bromide are similar. ,4s the total concentration of salt is increased, the dilute phase has increased salt and decreased sodium concentration; the concentrated phase exhibits the reverse behavior. In general, at a given temperature, adding salt permits lower and higher concentrations of sodium to remain in equilibrium with each other than would have been allowed in the absence of salt, thus broadening the phase separation curve. If the sodium metal data are plotted as in Figure 2 of
IJIQUID-LIQUID PHASE SEPARATION IN
ALKALIhfETAL-AMMONIA
a ~a
Br,-32.90'
X NOH,,
.08
0 Nal No Er,-S6*SV0 A NaI
0
3537
SOLUTIONS
NO NH,only, - 7 5 . 0 8
8 Na I, -71.00
3
NgI
No&
2
: -75.00:
0
-5639
0
-32.90' : -56,39' -32.9d
0
x
't
06
a
q 0.01
O0
0.01
'02
L
MOLE
Figure 1. Plot of mole fraction sodium in dilute phase os. concentrated phase. Data are given for three salts and three temperatures as Specified in the legend. At temperatures below the consolute point for sodium alone, points have been included for the consolute concentrations of sodium without added salt.
0
X a
Na8r,-32.90* NaNi
Nal
Figure 2 . Plot of mole fraction of sodium salt in dilute phase us. Concentrated phase. Data are given for three salts and three temperatures as specified in bhe legend. As is discussed in the text, the trend of data below the consolute temperature for sodium alone is entirely different from that above, i.e., at -33".
FRACTION
SODIUM METAL
Figure 3. Plot of mole fraction sodium salt us. mole fraction sodium metal for additions of sodium iodide, bromide, and azide as noted in the legend. Corresponding concentrated- and dilute-phase points should properly be connected by tie lines, but these are omitted for greater clarity. Refer to the text for a discussion of the appearance of these data on a suitably large-scale graph.
meiitary because of the low solubility of the salt, it is now apparent that the iodide behavior was not an artifact. Below the consolute temperature for sodium alone, the concentrated phase curves (as in Figure 3 ) proceed to intersect the sodium axis at a point corresponding to zero sodium salt, at the permitted sodium concentrations for the temperature in question, with the necessary consequence that all other points have larger concentrations of salt than zero, a concentration increasing with total salt concentration. Above the consolute temperature for sodium alone, the concentrated phase curve proceeds to a consolute point having concentrations of sodium and salt larger than any other value on the concentrated phase curve, which does not intersect the sodium axis, and which thus has a reverse slope to the concentrated phase curves at the two lower temperatures. At - 32.90°, the mole fractions of salt required to effect phase separation stand in the order YaI < S a S a < SaBr. The mole fractions of sodium at the extrapolated consolute points are the same, within 1.6%, approximately the precision of plotting and reading the graphs. In Figure 1, the data for all three salts fall on the same curve a t a given temperature, while in Figure 2 three separate curves result depending on the salt in question, thus demonstrating that while the effect depends upon the concentration of salt added, it depends
*
Schettler and P a t t e r ~ o n ,all ~ data will fall outside or above the temperature-composition curve, depending on the temperature. A concomitant effect of adding salt is to raise the consolute temperature above that permitted for metal alone. In the results reported earlier on sodium iodideJ2it was not clear why the behavior of the salt in the concentrated phase at -56 and -75" was the reverse of what has just been described. Even though the bromide results at -56" are frag-
(4) P. D. Schettler and A. Patterson, J . Phya. Chem., 68, 2865 (1964).
Volume 71, Number 11
October 1967
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PATRICIA WHITEDOUMAUX AND ANDREW PATTERSON, JR.
". N.1
Figure 4. Plot of mole fraction sodium vs. mole fraction sodium iodide (total). The three curves, reading from the top down, are sodium, total, with the actual data shown as the dotted curve, and a straight line drawn for comparison; middle curve, sodium, concentrated phase; bottom curve, sodium, dilute phase.
also on the salt used. Since each salt has the same cation, the distinctive effect must depend on the anion. Figure 4 is a plot of NNa us. NN,I(total). The three curves result from combining the data in ref 2 so as to yield ",(total), top curve; NNs(concentrated), middle curve; and .",(dilute), bottom curve, plotted against NN,I(total) in each case. This plot shows that total mole fraction of sodium is approximately linear in total mole fraction of added salt. The dilute and concentrated curves have been extrapolated to the sodiumonly point, on which they converge in a smoothly parabolic manner. Very similar curves result for azide and bromide though the range of salt concentration covered is smaller. The dependence of the effect on the salt used could arise in at least two different ways, a stoichiometric common-ion effect, or a specific ionic interaction. It can be shown, with data on two additional salts, that the first explanation is entirely unsatisfactory, the second more attractive. In the solutions studied here, the concentration of sodium lies (at temperatures above -78") from about 0.4 to 2.6 V I , dilute phase, and from 2.6 to 6 m,concentrated phase. Arnold and Patterson5 have proposed a model for these solutions which gives the species present in this range, in order of decreasing concentration, as S a + , Xaz, 31' (for example, S a e - or ea2-),e-, and Sao. The actual concentrations cannot safely be calculated, but the concentration of K a t is some 10z-fold greater than that of NaO, with the other species lying between. In assessing the possible interactions of these several species, the two neutral mes should play no important part, while the negative diamagnetic species would be less important in interactions with electrons since it is stabilized by spin pairing, though the negative charge T h e Journal of Physical Chemistry
might be important in electrostatic interactions. The solvation behavior of all the species might affect, the metal-salt interactions. Ionization constants have been determined for sodium iodide and bromide in liquid ammonia, though not for azide. The values are XaI, 2.80 X and NaBr, 2.898 X from the work of Kraus and his assoc i a t e ~ . ~ Gunn ~' and Green* have obtained a rather different value for sodium iodide through consideration of heats of solution of sodium in the presence of sodium iodide. The concentrations are nearer those in this work and iodide was present, but as there are no comparable dat'a for bromide, we must accept the data of Kraus6J and assume the azide behaves in a similar way. The small difference between the bromide and iodide ionization constants and the probability that the ionization of the salts is effectively suppressed by the large concentration of Naf present in the mixtures make these assumptions reasonable. It would be desirable to have data on sodium azide ionization; we propose to undertake a study to remedy this lack. In any case, the presence of an unionized species, S a x , is insufficient to account for the results through a common-ion effect. Catterall and Symonsg have studied esr g shifts in alkali metal-ammonia solutions in the presence of added salts. Sodium iodide gave a large negative Ag, bromide a small negative Ag, and chloride and amide a negligible shift. Azide was not studied. In the esr measurements, the effect of varying the alkali metal cation in the salt of a particular anion was noticeable only with iodide, but in any case the effect was much subsidiary to that of the anion. These authors considered that quantum mechanical interactions between solvated electrons and solvated anions or salt ion pairs were the important source of the effects they observed. The salt-metal ratios used were up to 200-fold larger than those employed in our study. The similarity of the results-the large effect of iodide and the smaller effect of bromidesuggests that a specific interaction between the solvated electron and an ion-paired salt is a much more likely source of the phase separation observations than a common-ion effect in which the original Braus equilibrium, Nao $ S a + e-, is shifted toward the left by added sodium salt. Indeed, the shift, owing to electron-salt interactions, is presumably to the right. As Figure 1 clearly shows. cation or metal behavior is not distinctively affected by the salt used,
+
(5) E. Arnold and A. Patterson, J . Chem. Phys., 41, 3089 (1964). (6) C.A. Kraus and W. C . Bray, J. A n . Chem. Soc., 35, 1315 (1913). (7) V. F. Hnizda and C . A. Kraus, ibid., 71, 1565 (1949). (8) S. R. Gunn and L. G. Green, J . Chem. Phys., 36, 368 (1962). (9) R. Catterall and M.C . R. Symons. J. Chem. Soc., 4342 (1964).
LIQUID-LIQUID PHASE SEPARATION IN ALKALIMETAL-AMMONIA SOLUTIONS
as is also found in ref 3, which shows the relatively minor effect of cation or metal interactions. It should be noted that, while the migration of the metal from the dilute to the concentrated phase with increasing added salt is a consequence of the effect of the salt in raising and broadening the temperature-concentration curve, the migration of the salt to the dilute phase is contrary to what would be expected from any of the commonly observed examples of phase separation. The recent work of Eisen and Jaffe'O is typical of the usual behavior: the concentration of the salt should increase in both phases. I n the present experiments, whether or not the temperature of observation is above or below the consolute point of sodium, the majority of the salt appears in the dilute phase. Just what role it plays there is a matter of interest and speculation. We have proposed in a separate paper" a model for two-component phase separation emphasizing the calculation of the chemical potential of the solute. Although the model is inadequate to cope with the threecomponent problem, in terms of the chemical potential, the effect of the salt is to raise it in the dilute phase and lower it in the concentrated phase, in spite of the lesser
3539
and greater concentrations of metal present. Taken together, the salt studies examined in this paper strongly suggest that electron-anion interactions make this possible, and that the order of the salts found depends on quantum electronic interactions with the anion, a possibility which should be checked by esr measurements on metal-azide solutions. The migration of the salt to the dilute phase is less easily understood. Since the metal is principally in the form of Na+, it may be that electrostatic effects are important. Alternatively, considerable quantities of solvent ammonia are bound to metal in the concentrated phase, so it may be that the dilute phase is a better solvent for the salt.
Acknowledgment. This work was supported by the National Science Foundation. The hospitality of the Physical Chemistry Laboratory, Oxford University, ext,ended to A. P. and the assistance and continuing interest of Dr. Paul Schettler during this work are gratefully noted. (10) E. 0. Eisen and J. Jaffe, J. Chem. Eng. Data, 11, 480 (1966). (11) P. D. Schettler, P. W. Dournaux, and A. Patterson, paper V in this series, submitted for publication.
Volume 71,Number 11 October 1967
