186
JOHN A. SKARULIS AND PAULKISSINGER
0.95 g.")than could a much smaller He atom in a hole of similar size.
Summary The lack of a great solubility difference between Ar and He in molten NaN03 leads to the conclusion that many of the dissolved gas atoms occupy holes of their own creation in the liquid, while other atoms may still occupy existing free volume. The Merence in the solubility effects on specific conductance may be interpreted in terms of (1) a greater dilution effect by Ar than by He, and (2) more effective blocking of ionic jumps by large Ar atoms than by smaller He atoms. As accurate density data for the gas saturated melt at various pressures become available, it conceivably
The Systems M,SiF,-(NH,),SiF,-H,O
may be found that the equivalent conductance isotherms corresponding to Figure 3 have more nearly the same slope. This, of course, would support the dilution mechanism as being the dominant effect.
Acknowledgments. The authors gratefully acknowledge support of this work by the National Science Foundation, Grant GP-4274. They are also grateful for interesting discussions with Dr. Milton Blander of North American Aviation, and with Dr. Ralph Seifert of Indiana University. This paper is based on the thesis of Walter C. Zybko, which has been submitted to the Graduate School of Kansas State University in partial fulfillment of the requirements for the degree of Doctor of Philosophy.
at 25" (M
=
Na, K, Rb, Cs)
by John A. Skarulis and Paul Kissinger Department of Chemistry, St. John's University, Jamaica, New York (Received October 1, 1966)
The systems M2SiFs(NH4)2SiF,H20 at 25" (M = Na, K, Rb, Cs) exhibit solid solutions and are Types IV, I, I, and V, respectively, according to the classification of Roozeboom. The distribution of ammonium ion between solid solutions and liquid solutions in the model systems M = K and Rb is described by the relationship xs = kxln in the ammonium hexafluorosilicate rich regions, xs and x1 being the mole fractions based on total moles of ions and water. A shift in the N-H frequency at 1420 em.-', observed in the infrared absorption spectra of solid solution phases except when M = Cs, is taken as evidence of hydrogen bonding which is postulated as being necessary for double-salt formation. The absence of the shift is believed to rule out a double salt suggested by the solubility data in the system M = Cs at 25".
and hydrogen halides, few with those derived from fluoro acids; all of the component hexafluorosilicate salts The data pertaining to the system Li2SiF6-(NH4)2in this series are anhydrous phases when in equilibrium SiF6-Hz0 at 25" and the double salt, LiNH4SiF6, with their saturated aqueous solutions, a situation have been reported.192 Presented here are the results rarely found in all of the alkali metal salts of a given of a study of the remainder of series of systems in which acid; the hexafluorosilicate ion is octahedral, a shape the alkali metal hexafluorosilicate was varied. This series is noteworthy for a number of reasons: most of (1) J. A. Skarulis, V. N. Darnowski, W. P. Kilroy, and T. Milazzo, the literature in the field of aqueous solubility equilibria J . Phys. Chem., 68, 3074 (1964). (2) R. Rudman and J. A. Skarulis, Acta Cryst., 18, 132 (1965). deals with salts which are derived from oxyacids
Introduction
The Journal of Physical Chemistry
THESYSTEMS M2SiFr(NH4)2SiF'rH20 AT 25" (M
= Na, K, Rb,
unlikely to be found among the anions of oxyacids; although hydrogen bonding apparently plays little or no role in the structure of pure ammonium hexaf l u o r ~ s i l i c a t ethis ,~~~ may not be so in double salts and solid solutions in which the latter is present; the application of infrared spectroscopy to diagnose hydrogen bonding of the N-H.-.F type in such new phases is possible since the ammonium ion and the hexafluorosilicate ion are both infrared active. The study of the systems of alkali metal hexafluorosilicates has been aimed a t providing additional experimental information pertaining to the criteria of isomorphism in salts and to the relationship between solid solution arid double-salt formation. I n connection with the latter, the essential role of hydrogen bonding in double-salt formation is generally recognized. Most of the familiar double salts are hydrated; those which are not hydrated generally contain the ammonium ion ;%sa structural unit. Evidently such bonds as O-H...O, O-H*..F, N-H.-.O, and NH . - + Fare necessary to their stability. In the systems reported here, only N - H - . * Fbonds can have an effect upon the structure of any new solid phase since the phase is likely to be anhydrous. A tentative hypothesis in all this has been that anhydrous double salts in which the lattice units are spherically symmetrical ions probably do not exist. In such cases, solid solutions are possible. To form double salts, in any case, there must be a directed character to the ionic attractions, most often provided by hydrogen bonds which assume tetrahedral symmetry. To cite an example, the systems NHZ-(NH4)2SiF~H20at 25" (X = F, GI, Br, I), in which all solid phases are anhydrous, were found to be simple except when X = F.586In this case, where the strongest N-H...X bonds exist, the double salt, NH4F- (NH4)2SiF6, occurs as a phase. However, if a salt, in which hydrogen bonding is not possible or is weak, should have a formula which corresponds to a double salt, such as AXaBX, then it is suggested that the salt is actually a complex salt, a lattice of two ions, rather than a double salt, a lattice of three ions. The kainites, M1C1.MI1SO4, double salts with apparently four ions, may be cited as a similar case. To account for the observation that molten KCI-MgS04 exhibits a minimum electrical conductance in comparison to mixtures in which the ratios of the component salts is other than 1:1, the explanation has been offered that K+ and CIMgS04ions exist in liquid kainite and, presumably, in the solid.' From a statistical viewpoint, a regular threedimensional lattice of four different ions, as in a heteroionic double salt, is highly improbable; one of only
Cs)
187
three ions is still highly improbable unless bridging units are present.
Experimental Section The equipment and general procedure were the same as described in the system involving lithium hexafluorosilicate. The ammonium hexafluorosilicate (99.9%) was also the same except that it was heated a t 110" prior to use. The alkali metal hexafluorosilicates were prepared from C.P. or reagent grade metal chlorides and 30% hexafluorosilicic acid. The rubidium and cesium chlorides were obtained from Fielding Chemical Co., and the others from Baker Chemical Co. In a typical preparation, 150 ml. of a filtered solution of chloride (0.40 mole) was added dropwise with stirring and cooling to 100 ml. of 30% acid (0.25 mole) which also had been filtered after a preliminary chilling. The mixture was then stirred for at least 0.5 hr. without any further cooling. The precipitate was washed exhaustively by decantation with water acidified with hexafluorosilicic acid, then transferred to a Buchner funnel, filtered, and washed repeatedly until the filtrate gave a negative test for chloride ion. After being washed with 95% ethanol, it was air-dried for several days and finally heated at 110" prior to use. The hexafluorosilicate analysis of the final products, by titration in hot solution with standard sodium hydroxide and phenolphthalein as indicator, gave assays of 99.6, 100.0, 100.7, and 100.4% for sodium, potassium, rubidium, and cesium hexafluorosilicates, respectively . Known mixtures of the component salts and water were equilibrated over periods ranging from 3 to 6 weeks a t 25 f 0.02" by means of an end-over-end motion in polyethylene test tubes which contained three Teflon balls for more efficient mixing. The saturated solutions were analyzed after various time intervals of rotation for total alkali metal, ammonium, and hexafluorosilicate ions by the cation-exchange procedure described in the work mentioned.' Constancy of composition within the limits of precision of the method of analysis (5 parts per 1000) was taken as an indication that equilibrium had been attained. The final analysis of a given saturated solution included a separate determination of alkali metal ion. With the exception of those (3) B. Cox and A. G . Sharpe, J. Chem. SOC.,1798 (1954). (4) C. C. Stephenson, A. Wulff, and 0. R. Lundell, J . Chem. Phys., 40, 967 (1964). (5) J. E. Ricci and J. A. Skarulis, J . Am. Chem. SOC.,73, 3618 (1951). (6) J. A. Skarulis, R. E. Connolly, and W. M. Wypyski, unpublished data on X = Br, I. (7) N. P. Luzhnaya, Dokl. A M . Nauk SSSR, 69, 809 (1949).
Volume 70, Number 1 Januarw 1966
188
JOHN A. SKARULIS AND PAUL KISSINGER
containing sodium, samples were evaporated to dryness, ignited at 400" to volatilixe the ammonium salt, and the residues finally weighed as hexafluorosilicates. Samples in the sodium system were ignited at 600' instead, and the residues were weighed as sodium fluoride. In volatilizing the ammonium salt from various samples, particularly those last mentioned, an iridescent film, presumably silicon dioxide, formed on the crucible and cover. As a precaution, 10 drops of 50% hydrofluoric acid were added to the crucible and evaporated. This served to remove the film. Except in the case of sodium fluoride residues, this treatment was followed by one with an equal number of drops of 30% hydrofluorosilicic acid before a final ignition at the temperature indicated. As a check, the weighed hexafluorosilicate residues were titrated hot directly in the crucibles with sodium hydroxide, the smaller residues with 0.01 N base. I n addition to the measurements just described, the solubilities of rubidium and cesium hexafluorosilicates in pure water were checked. Samples of saturated solutions were evaporated to dryness, the residues dried at 110" and weighed as hexafluorosilicate. Also, a number of solid phases in the ternary systems were subjected to infrared examination. I n preparation for this, samples were filtered, washed with 95% ethanol, and air-dried. The spectra were observed with a Perkin-Elmer Infracord spectrophotometer on KBr disks which contained 0.2-0.4% by weight of hexafluorosilicate mixture. These were compared with the spectra of the component salts and lithium ammonium hexafluorosilicate for evidence of hydrogen bonding, on the assumption the N-H * * F bonds were present in the double salt and were necessary to its stability.
HzO
(W)
A
(NHdzSiFs (B)
NazSiFe (A)
Figure 1. The system NasSiFs-(NHI)2SiFG-H20 a t 25'. HzO (W)
-
Results and Discussion The results of the analyses of the saturated solutions are summarized in Table I. Except when M = Na, the absolute error of the percentages of alkali metal hexafluorosilicatein the saturated solutions is 0.0295, an error inherent in the weighing of small residues in the analytical procedure. The weights of hexafluorosilicate residues which were calculated from the titrations with base agreed within this range with those obtained by direct weighing. When M = Na, the estimated error is 3=0.05%. The data are plotted in Figures 1-4. There is evidence of solid solutions in all four systems, the solid solutions being continuous when M = K or Rb, but exhibiting a gap in miscibility when M = Na or Cs. Included in the table are the mathematically extrapolated compositions of the solid solutions on an anhydrous basis. Although complete analyses of these solid phases were not made, drying
*
The Journal 0.f Physical Chemistry
KlSiFe (A)
(NHdzSiFe (B)
Figure 2. The system K&3iF6-(NH4)2SiF6-H20 at 25".
tests on selected ones were made to establish that they were actually anhydrous. According to the classifkation of Roozeboom,*the system is Type I when M = K or Rb, Type IV when M = Na, and Type V when M = Cs. The formation of continuous solid solutions between ammonium hexafluorosilicate and the potassium and rubidium salts is not unexpected Moresince a t 25" all three have cubic ~ t r u c t u r e . ~ (8) B. Roozeboom, 2. physik. Chem., 8 , 521 (1891).
THESYSTEMS MkliF~-(NH~)tSiF,j-H20 AT 25' (M = Na, K, Rb, Cs)
189
~~~~
Table I: Systems MzSiF6 (A)-(NH&SiFe (B)-HzO (W) a t 25" Solid
-Complex%A %B
-Solution%A
%B
M
... 9.23 8.93 8.85 9.20 9.05 9.28 2.97 2.18 4.90 2.93 1.57
RbzSiFa (A)
(NHdzSiFs (B)
Figure 3. The system Rb2SiFr(NH4)2SiFsHz0 at 25". Hi0 (W)
... 5.05 9.88 15.32 17.23 18.02 19.02 20.56 19.27 24.21 25.09 23.28
0.81" 0.96 1.15 1.27 1.46 1.18 1.32 1.37 1.38 1.29 1.41 1.51 Av. 1.35
...
. ..
10.06 9.65 10.00 9.82 7.01 3.58 1.42 0.52
1.85 7.44 12.12 17.50 18.02 20.66 23.52 23.05
...
...
9.99 10.10 10.17 10.04 7.13 4.86 2.99 0.98
1.98 6.09 9.98 17.95 21.09 23.07 25.06 27.12
...
...
(NH4)zSiFa(B)
O Figure 4. The system C S Z S ~ F ~ ( N H & ~ F ~ - aHt Z25".
over, the molecular volumes of the latter two, calculated from lattice dimensions, are 135.0 and 150.4 X cm.a, respectively, which differ by only 9% from 148.0 X cm.3, the molecular volume of the former. In contrast to these, although cesium hexafluorosilicate is also isomorphous with the ammonium hexafluorosilicate, its molecular volume of 175.0 X 10-24 cm.a is 18% higher. In this case, complete miscibility is less likely and this expectation is borne out by the experimental data.
=
7 N H 4+ZB
XI
Na
0.00
5.36 10.49 16.38 18.23
0.00 1.71 2.97 2.99 6.31
0.000 0.012 0.021 0.021 0.044
0.0000 0.0113 0.0231 0.0378 0.0428
l8.,ll 18.70 18.591
i:
linvariant I
18.76 18.48 18.63)
0.13' 0.09 0.13 0.13 0.12 0.06 0.04 0.01 0.06
M = K 0.00 1.78 4.49 6.72 9.33 10.61 13.78 16.97 18.21
0.00 2.46 27.1 39.7 50.7 56.7 70.7 85.3 92.9
0.000 0.020 0.209 0.299 0.373 0.411 0.499 0.585 0.627
0,0000 0.0037 0.0094 0.0143 0.0203 0.0232 0.0309 0.0390 0.0421
0.215" 0.15 0.11 0.13 0.14 0.16 0.09 0.07 0.03
M = Rb 0.00 1.63 4.59 6.90 10.81 12.96 14.56 15.99 17.75
0.00 5.00 17.0 28.7 48.1 59.7 69.3 79.5 92.4
0.000 0.057 0.177 0.277 0.412 0.481 0.532 0.581 0.637
0.0000 0.0034 0.0096 0.0147 0.0237 0.0290 0.0328 0.0364 0.0410
o.953c 0.95 1.32 1.32 1.34 1.33 1.33 1.27 1.27 1.20 1.09 1.05 1.04 0.72 0.38 0.00
M = cs 0.00 3.19 9.70
0.00 0.0000 0.000 0.52 0.0079 0.0067 2.58 0.038 0,0213 0.185 0.0212 0.253 0.0212 9.65 invariant 0.323 0.0212 Av. 9.66 9.89 13.03 10.36 31.3 0,340 0,0229 10.15 13.93 10.99 32.8 0.352 0.0244 10.02 14.99 11.86 34.6 0.365 0.0265 9.90 16.36 12.89 37.2 0.384 0.0289 10.02 17.35 13.57 38.9 0.395 0,0306 10.11 18.25 13.96 41.3 0.411 0.0316 17.10 65.8 0.543 0.0396 4.89 23.17 0.98 27.12 18.73 94.2 0.649 0,0437 ... ... 18.75d 100.0 0.667 0,0437 Citing recent solubility values only: interpolated from values a t 10 and 50' of N. S. Nikolaev, N. -4. Ivanov, and S. G. Koltypin, J . A p p l . Chem. USSR, 9,1183(1936). ' Interpolated from values at 20 and 40" of I. G. Ryss, Russ. J . Phys. Chem., 21, 197 (1947). This work. See ref. 5.
10.18 2.94 9.96 9.06 10.02 10.14 10.15 10.96 9.96 12.05 CszSiFs (A)
solution, %B
Volum 70, Number 1 January 1966
JOHN A. SKARULIS AND PAUL KISSINGER
190
However, the sodium hexafluorosilicate system is more complicated. Unlike the others, the sodium salt has a hexagonal structure at 25019but this does not rule out solid solutions since the pair of salts need not have the same crystal structure.1° Two conclusions are possible from the solubility data alone: either the solid solutions are continuous, Rooeeboom Type 11, or they are discontinuous, Type IV. The sole support for the first conclusion is that the solubility curve in Figure 1 bulges away slightly from the 0% NazSiFB edge of the triangle and, therefore, would seem to pass through the point of minimum water content in a continuous manner. However, since the curve is so close to the edge, the minimum point could be equally the intersection of two curves, thus making the system Type IV. This second alternative can be supported by various plausible arguments. (1) If sodium hexafluorosilicate were cubic at 25", its molecular volume would be considerably different from that of the ammonium salt so that at best, by analogy to the cesium case, only partial miscibility could be expected, ruling out Type 11. ( 2 ) Assuming a miscibility gap, it is obvious from Figure 1 that solid ammonium hexafluorosilicate is much more soluble in solid sodium hexafluorosilicate than sodium hexafluorosilicate is in ammonium hexafluorosilicate. This is to be expected from the following considerations. When the ammonium salt is at 25", it is at a point not too far removed from 14") the temperature of its hexagonal cubic transition.1' Thus, it is not too far from the point a t which the hexagonal lattice is the stable one. It is reasonable to presume that the ammonium ion can replace the sodium in a hexagonal lattice a t 25" up to a point a t which the distortion of the lattice is so great that no further substitution can take place. This may be viewed a i a case of the limited solubility of metastable hexagonal ammonium hexafluorosilicate in stable hexagonal sodium hexafluorosilicate a t 25". In a similar manner, the low solubility or, perhaps better, the insolubility of the sodium salt in the ammonium salt is understandable. The cubic form of sodium hexafluorosilicate, if it did exist, would be stable only at temperatures far removed from 250.12 Therefore, it is highly improbable that the sodium ion, which prefers the hexagonal arrangement strongly, will substitute for ammonium ion in the cubic form of ammonium hexafluorosilicate. The manner in which components distribute themselves between solid solution and liquid solution i c , of interest. I n the case of a pair of alums which form a continuous series of solid solutions, the distribution has been described by the empirical equation The Journal of Physical Chemistry
R1
=
KR,"
in which R1 and R,represent the mole ratios of the two positive ions by which the pair differ, K is the true distribution constant in terms of activities, and m is a constant related to the attraction and repulsion of the components in the solid phase.13 Equations of this type cannot be applied to the data in Table I because of the large relative errors in the mole ratios of ammonium ion to alkali metal ions. However, the data are sufliciently accurate to correlate the distribution of ammonium ion only. To this end, the mole fractions of ammonium ion, x, and 51, based on
7.0
t-
I 0.0
1.0
2.0 ZI
x
3.0 10%.
4.0
6.0
Figure 5 . Distribution of NH4+ ion in the systems M&XFe-(NH&SiF6-Hz0 at 25" (M = K, Rb, Cs): zsis the mole fraction in the solid phase; X I is the mole fraction in the liquid phase, based on the total moles of ions and water. (9) B. Cox, J . Chem. SOC.,3251 (1954). (10) A. F. Wells, "Structural Inorganic Chemistry," 3rd Ed., Oxford University Press, London, 1962, p. 185. (11) I. G. Ryss, "The Chemistry of Fluorine and Its Inorganic Compounds," Translation Series, U. S. Atomic Energy Commission, Oak Ridge, Tenn., 1960. (12) E. S. Freeman and V. D. Hogan, A d . Chem., 36, 2337 (1964). (13) A. E. Hill, G. S. Durham, and J. E. Ricci, J . Am. Chem. SOC., 62, 2723 (1940).
THESYSTEMS RI12SiF,-(NH+)zSiF6-Hz0 AT 25" (M
Log -2.5
K, Rb, Cs)
191
2.1.
- 1.5
-2.0
= Na,
- 1.0
-I
O K
P
/
@
Rb
9
cs
I
1 Cm.-1.
3
Figure 7. Infrared absorption spectra in KBr of solid phases containing ammonium hexafluorosilicate: A, pure ammonium hexafluorosilicate; B, lithium ammonium hexafluorosilicate; C, D, and E, solid solutions of ammonium hexafluorosilicateand potassium, rubidium, and cesium hexafluorosilicates, respectively.
- -2.0 -1
!
Figure 6. Distribution of N&+ ion in the system MzSiF&N€I&XFsHzO a t 25" (M = K, Rb, Cs): z8 is the mole fraction in the solid phase; z1 is the mole fraction in the liquid phase, based on total moles of ions and water; when M = Cs, only the continuous solid solutions on the ammonium hexafluorosilicate rich side of the miscibility gap are represented.
total moles of ions and water, in the solid and liquid phase, respectively, are tabulated in Table I and plotted in Figure 5. Numerical analysis of these values indicates that the expression xs
=
kxln
(2)
is applicable over a considerable range of concentrs tions in the systems M = K and M = Rb. In the form shown, eq. 2 resembles the familiar Freundlich adsorption isotherm which has been applied to the distribution of a solute between a solid adsorbing phase and a solution. Taking the common logarithms of both sides of eq. 2 leads to Log xs = log k n log X I (3)
+
which indicates that a graphical plot of log xs us. log 21 should be a straight line with a slope of n. Pertinent parts of the data of three systems are so plotted in Figure 6, excluding M = Na, with attention focused upon the ammonium hexafluorosilicate rich portions. Seven of eight experimental points when M = K, representing 75% of the solubility curve, fall on a straight line which has a slope of 0.70. The eighth point is obviously in error and is omitted. When ,nil = Rb, the points fall on a smooth curve which
approaches this line as a limit. However, when M = Cs, the points on the ammonium hexafluorosilicate rich side of the miscibility gap fall on a curve whose limiting slope is 2.0. The systems M = K and M = Rb may be considered as being model systems for testing eq. 2 or any other distribution relationship. It has been noted above that the molecular volumes of potassium and rubidium hexafiuorosilicates differ by 9% from that of the ammonium salt; one is smaller, the other larger. Therefore, the data of these systems show the effect of substituting a cation which produces either a small negative distortion or a small positive distortion in the vdume of the ammonium salt. It would seem from this that eq. 2 should hold as long as the distortion is small. Also, it is noted that the solubilities of the two alkali metal salts, given in Table I, are almost the same. This means that the sum of the lattice energy and the hydration energy terms of the distribution process a t a given concentration of solution is about the same for both. Similar behavior is expected from both systems because of the above reasons. Apparently, the system M = Rb is less model than the system M = K. If size only is considered, the deviation of the system M = Cs is expected because substitution of the considerably larger cesium ion produces a distortion which is too great. IC this connection, it is evident from Figure 5 that this substitution reaches a limit a t xs = 0.33, or a t a mole ratio of cations of 1:1, beyond which Volume 70,Number 1
January 1966
192
there is a break in miscibility. Since the tie lines of this system in Figure 4 show no tendency to converge a t this point, a conclusion that this limiting ratio represents compound formation cannot be drawn. A study of this system a t other temperatures would serve to establish this point. The solubilities of rubidium and cesium hexafluorosilicates which were determined in this study appear in Table I. They agree well with recent values in the literature.14 However, only an approximate comparison of values is possible because these recent solubilities If one assumes are expressed in units of moles lite+. the density of the solutions to be 1.00 g. mI.-l, the converted values of the solubilities of the two salts are 0.216 and 0.96% by weight, respectively. The infrared absorption curves of pure ammonium hexafluorosilicate and homogeneous solid phases containing it are given in Figure 7. The latter samples were selected so that the mole fraction of ammonium hexafluorosilicate in each was about the same (0.5) except for the rubidium hexafluorosilicate solid solution (0.2). Curve D, which is typical of the alkali metal hexafluorosilicate rich phases in the potassium and rubidium systems, is presented because it shows a curious sharp rise in transmittance around 1350 em. -l for which no explanation is offered here. I n general, the solid phases of these two systems which were examined gave similar spectra. Considering spectrum A, the band with a peak a t 748 em.-' is characteristic of the hexafluorosilicate ion; the ones with peaks a t 1420 and 3020 em. -1 are characteristic of the ammonium ion. The range 1390-1420 em. -I, assigned to the N-H vibration frequency in ammonium salts of chloro and fluor0 complex acids, has received special attention. A shifting beyond this range has been taken as an indication of hydrogen bonding.a Such a shift, in which the shoulder a t 1440 cm.-l just perceptible in curve A becomes relatively more prominent, is evident in B, C, and D. It is discernible also in a spectrum of a solid phase of the sodium hexduorosilicate system (whose spectrum is not shown in Figure
The .Tournal of Physical Chemistry
JOHN A. SKARULIS AND PAUL KISSINGER
7 but which has the features of curve B) taken from the sodium hexafluorosilicate rich side of the miscibility gap represented in Figure 1. The shift is accompanied by one a t the second characteristic frequency of the ammonium ion, in which the shoulder a t 3035 em.-' becomes more prominent, and by a broadening and shifting of the Si-b band a t 748 cm.-l. On the basis bonds play a of this, it is concluded that N-H.*.F significant role in hhe structures of lithium ammonium hexafluorosilicate and the solid solutions studied except those involving cesium hexafluorosilicate. Referring to this last case, curve E differs from curve A by only a slight broadening of the band a t 3020 em.-' which could be the effect of traces of moisture. Though the region around 3020 em.-' may be more sensitive in diagnosing N-H * F bonding, the occurrence of N-H. .O and O-H. * -0 frequencies in the near vicinity makes a reliable interpretation difficult. A calculation making use of crystal and bond distance data can be made to show that the N-H-F distance in ammonium hexafluorosilicate is just a t the upper limit for stable hydrogen bonding. Introduction of a smaller ion into the lattice can produce distortion which will bring certain of the N-H-F units close enough for hydrogen bonding. However, introduction of a much larger ion such as cesium can only increase the lattice dimensions and make hydrogen bonding less likely. Thus, there is reason to believe that the 1 : l limiting ratio of the solubility of cesium hexafluorosilicate in ammonium hexafluorosilicate does not represent the formation of a double salt. I n conclusion, various combinations of pairs of alkali metal hexafluorosilicates and water are under investigation. These results should be of particular significance since no hydrogen bonding is possible in the solid phases if they prove to be anhydrous as expected. In addition to infrared absorption data, Xray data on all these systems are being gathered. (14)R. H.Schmitt, E. L. Grove, and R. D. Brown, J. Am. Chem. SOC.,82, 6292 (1960).
