J. Phys. Chem. 1982, 86, 4351-4356
paramagnetic species, and its wealth of information concerning the structure and kinetics of these radical ions, constitutes ,a unique tool for elucidating the behavior of radical ion ripecies in pulse radiolysis. Acknowledgment. We acknowledge Professor Ronald Lawler for several useful discussions and Dr. Marion
4351
Thurnauer for her assistance with the cryogenic ESR experiments. We also thank Rober Lowers and Alan Young for technical assistance and operation of the Van de Graaff accelerator. This work was performed under the auspices of the Office of Basic Energy Sciences, Division of Chemical Science, U.S.Department of Energy under Contract No. W-31-109-ENG-38.
Solubilitic3s of Small Molecules in Liquid Krypton W. H. Beattie,' Wllilam B. Maler 11, S. M. Freund, and R. F. Holland Universw of California, Los Alamos Natlonal Laboratoty, Los Alamos, New Mexico 87545 (Received: May 10, 1982; In Final Form: Juiy 6, 1982)
Solubilities of MoFs, C02,CH20,HN,,and CH30H in liquid krypton have been measured over temperature intervals between 118 and 165 K. The solubilities at 140 K vary from = 0.1 mol/L for C02 to =2 X lo-, mol/L for HN,. Relative solubilities of CCl, and SO2 in liquid krypton have also been determined. The solubilities are compared with the Hildebrand-Preston-Prausnitzformalism of regular solution theory. Methanol is found to dissolve in liquid krypton as dimers or higher aggregates. Integrated cross sections have been measured for a few absorption features of the dissolved compounds.
I. Introduction Studies of spectroscopy and photochemistry of cryogenic solutions of liquified gases14 have shown them to possess several useful properties. Infrared spectra are simplified, and spectral. peak heights are enhanced by the collapse of the rotational band structure. Reactive compounds or mixtures m,ay be stabilized a t relatively high molecular densities. In all spectral regions, there are solvent-induced spectral shifts. Liquified rare gases are transparent in the commonly used spectral regions, chemically unreactive, and are relatively good solvents for many compounds. The solubilities of compounds in the liquified rare gases are therefore of interest. In this wolrk, solubilities of several compounds in liquid krypton have been measured between 118 and 165 K. Heats of sollution have been calculated, and the experimental data are compared with predictions from ideal and regular solution t h e o r i e ~ . ~ ? ~ We first present the formalism associated with regular solution theory. The expressions obtained from tlis theory for heat of solution and solubility of solid solutes in liquids are given. For an ideal solution, the heat of solution is equal to the heat of fusion. In section IV, the measured solubilities ,and heats of solution are compared with the (1) W. B. Maier 11, S. M. Freund, R. F. Holland, and W. H. Beattie, J. Chem. Phys., 69,1961 (1978). (2) S. M.Freund, W. B. Maier 11, R. F. Holland, and W. H. Beattie, Anal. Chem., 5;0, 1260 (1978). (3) S.M. Freund, W. B. Maier 11, R. F. Holland, and W. H. Beattie, J. Am. Chem. ,Sot., 101,4522 (1979). (4) W. H. Beattie, W. B. Maier 11, R. F. Holland, S. M. Freund, and B. Stewart, "Proceedingsof the 22nd Annual Symposium of the Society of PhoteOptiud Engineers, San Diego, CA, Aug 31,1978", SPIE Vol. 158, Laser Spectroscopy. (5) J. H. Hildebrand, J. M. Prausnitz, and R. L. Scott, "Regular and Related Solutions",Van Nostrand-Reinhold,New York, 1970, pp 22,87, 99 ff, 207 ff. (6) J. M. Prausnitz, "Molecular Thermodynamics of Fluid-Phase Equilibria",Prentice-Hall, Englewood Cliffs, NJ, 1969 pp 316 ff, 388 ff. 0022-3654/82/2086-4351$01.25/0
theoretical values, and the applicability of regular solution theory to the systems investigated in this work is discussed. In a real solution, the solute activity, u2,a t a temperature, T, is given by5v6
(1) where AH: is the heat of fusion of the solute at its freezing temperature, To. R = 1.988 cal/(mol K). AC, = C: - C," difference between the heat capacities of the liquid and solid solute. AC, is assumed to be independent of T. It can normally be anticipated that AC, 5 3R/2 =3 cal/(mol K), because translational modes are partially frozen out in the solid. Since, moreover, T - Tohas opposite sign from In (T,/T), the term AC,[T - To + T In (To/T)] is small and will be neglected here. (In this study, the term is largest for MoFs, where To = 291 K. With T = 140 K, this term is =72 cal/mol.) The expression5v7 R T In (u2/X2) = U Z $ J ~ ~ [ (-~ I82)' + 24261621 (2) from regular solution theory has been used to describe the activity coefficient, u2/Xz,of nonideal liquid solutions of solid solutes.g11 41 = XlUl/(XlUl + XZUZ) (3) u1 and u2 are respectively the molar volumes of the solvent and of the supercooled liquid solute. XI and X2 are re(7) G.T. Preston and J. M. Prausnitz, Ind. Eng. Chem. Process Des. Deu., 9, 264 (1970). (8) G. T. Preston, E. W. Funk, and J. M. Prausnitz, J. Phys. Chem., 75,2345 (1971). (9) E. Szczepaniec-Cieciak, B. Dabrowska, and J. M. Lagan, Cryogenics, 17,621 (1977). (10)Z. Woitaszek and E. Szczepaniec-Cieciak, Cryogenics, 15, 257 . (1975). (11) Z.Wojtaszek, E.Szczepaniec-Cieciak,and A. Morzyniec, Cryogenics, 15, 351 (1975).
0 1982 American Chemical Society
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cell is then filled with liquid krypton by condensing krypton vapor and is held a t a chosen temperature. The mixture is stirred vigorously for periods ranging from 15 min to 2 h to establish equilibrium between the solid solute and the saturated solution.lZ The integrated absorbance, A , of one or more infrared absorption features characteristic of the solute is measured with a Perkin-Elmer Model 180 double-beam spectrophotometer. The molecular number density, n, of the dissolved solute is then calculated from
H I cm
A = Qnl
C' Figure 1. Cryogenic cell for measurement of absorption cross sections. The cell body is copper. a and b are respectively fill and coolant pipes; c is a channel in the cell body. The windows, e, are AgCl and are held on by screws and spring washers, d, and clamping plates, f. g is a resistor. h is a recess to reduce thermal conductivity between the upper and lower parts of the cell. The optical path length through the cell is 1.3 cm, and the volume of the cell Is 2.6 cm3. Thermocouples are attached near the top and bottom. A Teflon-coated magnetic stirring bar rests on the bottom of the cell. The cell is held in a vacuum jacket having KRS-5 windows.
spectively the mole fractions of solvent and solute in solution. & and 6 , are respectively the solubility parameters of the solvent and solute; they can be estimated from 6 , = [(Mi" - R7')/~i]'l~ (4) where AH? is the molar heat of vaporization of compound i (=1or 2). 112 is a number which is characteristic of the given solute-solvent pair.7 The value of 112 is not wellknown but is predicted to be and nearly independent of temperature and composition.8 Equation 2 differs from the usual Hildebrand-Scatchard equation5 in that it contains the term 211d16,; effectively 112 is sometimes treated as an adjustable parameter whose value is determined by the characteristics of the solution." The heat of solution, AH0, is given by d In X 2 / d T = AH'/(RF) The integrated form of this equation is In X 2 = K - AHO/RT
(5) where K is a constant of integration. To obtain K and AH8 from regular solution theory, one combines eq 1 and 2 and puts the result into the form of eq 5 . Then AH' = AH,f + V Z ~ ~ ~ [-( 62)' S , + 24,61621 (6) for the effective heat of solution, provided ACp = 0 and
K = AHof/RTo
(7)
11. Experimental Methods
Solubilities are determined as follows. A known amount of solute is condensed into the cell shown in Figure 1. The
(8)
where 1 is the optical path length, 1.3 cm, through the cell and Q = .fa(?) d? is the integrated absorption cross section (in cm2 cm-') of the infrared absorption feature. a@)is the absorption cross section (in cm2)of the chosen infrared band, and 3 is the wavenumber (in cm-l). This procedure requires that Q be independent of temperature, T. Experimental tests for this requirement have been confined to the MoF, and COz absorption features in this work, and in these cases Q seems to be sufficiently independent of T. In any case, one does not expect Q to depend strongly on T because the density of liquid krypton does not change greatly over our temperature range and because the solute molecules reside primarily in their ground vibrational levels a t these temperatures. The following materials were used without further purification: Kr (Air Chemicals and Products, research grade, 99.995% pure), COz (Matheson, 99.5% pure), SO2 (Matheson, 99.98% pure), MoF, (Ozark-Mahoney), CH,OH and CC14 (Mallinckrodt AR grades). CHzO monomer was prepared as described in ref 1 and stored at liquid-N2 temperature. Hydroazoic acid, HN3, was prepared in a glass system equipped with Viton 0-ring-sealed valves by mixing 0.6 g of NaN3 (MCB, practical grade) with 5.5 g of stearic acid (USP grade), by evacuating to remove air, and by heating slowly to 70-90 OC; HN3 vapor was collected and stored at room temperature and a t pressures of C50 torr to avoid spontaneous explosions. The interior of the cell in Figure 1 is connected to a vacuum system through the fill pipe, a. The cell is cooled to the desired temperature by forcing timed bursts of liquid nitrogen through the coolant pipes, b, and channel c. The upper part of the cell can be heated by passing current through the resistors; a temperature differential up to 20 "C can be maintained across the thin-walled section of the cell a t h to prevent solute from depositing in the upper part of the cell. The minimum temperature differential in the cell body between points near g and c in Figure 1is 7 OC at -158 "C and 3 "C a t -110 "C. Cell temperature is controlled from the bottom thermocouple. Temperature excursions of the cell are CkO.3 OC. Usually, the amount of solute to be put into the cell was determined by measuring the pressure of solute vapor in a calibrated volume with a MKS Baratron capacitance manometer. A temperature differential was established between the top and the bottom of the cell, and the solute was condensed into the bottom of the cell. Any residual solute was flushed-into the cell with krypton gas. The amount of solute actually in the cell should be known to within +5% to -15%, maximum uncertainity. The larger negative error results from possible losses during filling. Different procedures were required for HN3, because it is unstable, and for CC4, because it has a low vapor pressure at room temperature and a high solubility in (12) Throughout the temperature range covered, the temperature of the mixture was set sometimes by cooling the cell and sometimes by warming it.
The Journal of Physical Chemistry, Vol. 86, No. 22, 1982 4353
Solubillties of Small Molecules in Liquld Krypton
TABLE I: Summary of Measurements molecule
lo-l%l,o,a
A P , b
molecules/cm3
kcal/mol
1O1aQ,C cm2 cm-'/molecule
vibrational moded
cm-l
2.87 k 0.1 1345-14105 1.1 v2 + v 3 1.868 i 0.08 20.3Bh 2 v 1 t 2va, u , + v , + v , 1500-1630 3.06 * 0.16 6.4' v2 635-65V 4.47 f 0.2 1075-1250k 26.1h v1 so2 5.70 0.4 16' 2100-2 148 HN, CH,O 1 6 50-1825 5.09 i 0.2 6.5 VZ CH,OH 980-1 1 6 0 6.19 f 0.2 20 v 7 , u s , (Vll?) Solubility a t 1 4 0 K. See e q 9. Parameter defined in eq 9 and approximately equal to AHS. The quoted errors are taken from the uncertainties in the slopes of the lines in Figures 2 and 3. Integrated cross section for the absorbance features lying within the spectral range A 7 given in the last column. Derived from eq 8. Molecular bands thought t o contribute t o &. e A T = spectral range over which absorptions were integrated t o obtain A in eq 8. f The absorbances of MoF, between 890 and 940 cm-' and between 1450 and 1500 cm-' were also utilized. 8 Based o n a crude estimate + 5 0 % ) of the CCl, in the cell and the fact that apparently n o t all of the CCl, dissolved a t the highest temperature reached. J;Lower limit o n Q,.' Computed from A for the "CO, v 2 band at 1.11%natural abundance of 13C,and the known amount of CO, in the cell. Because of the high absorption cross sections for the CO, features, the v , band of W O , was used. At low concentrations, the integrated absorbance between 1 2 7 5 and 1425 cm-' was used. This value is more uncertain than the others in this column because of possible decomposition of HN, and because of the procedure for getting HN, into the cell. MoF, CCl, COZ
6.8 167g 64 L0.54 0.14' 6.4 0.26
_+
'
J
IO0
I
I
-
1
IO
I
m
E
0 \
v, W
0"W
1.0
-I
0
I a,
-0
R E L A T I V E SOLUBILITY IN L K r
*
k
z"
0.1
W
n U
I
W
m
I 2
z 0.0
0.001
Flgure 2. Solubilities of GO,, MoF,, CH,O, CH,OH, and HN, in liquid The krypton (LKr) vs. 1 / T . Data are indicated by X, 0 , A,H, and parameters n and AHm ghren in Table I are obtained from the solid lines. Where given, the dashed lines indicate the concentration corresponding to the total amount of solute used in the experiment.
+.
liquid krypton. In both cases, the top of the cell was warmer than the bottom when the solute was added. The HN3 was first mixed with 1 atm of krypton gas and was then rapidly flushed into the evacuated cold cell with a higher pressure of krypton gas; the amount of HN3actually
l
l
,
,
I
,
The Journal of Physical Chemistry, Vol.
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86,No. 22, 1982
Beattie et al.
TABLE 11: Solubility Parameters, Heats of Solution, and Related Data aHOf+ molecule
To,aK
u , ~ cm3/mol
MoF, CCI, COZ SO, CH,OH krypton
290.7 250.3 216.6 197.7 175.5 116.0
67.98 82.30 33.60 35.12 33.70 35.82
aHof,d
cm-3'2
calimol
AH^,^ cal/mol
8.834 9.140 13.08 12.72 15.58 7.252
2198 597 2057 1768 758
3023 2046 3180 4620 6343
6,c
L'Z(6 i SZY,
calimol 2368 890 3198 2818 3095
10-18n, l o , molecule/cm3 112
calcdg
obsdh I
0.075 0.106 -0.003 0.255 0.427
14.4 36 22 0.093 1.9 x 10-5
6.8 267 64 2 0.54 0.26
Melting point, taken from ref 13. Molar volume at 1 4 0 K of the liquid krypton or supercooled liquid solute, inferred from liquid densities in ref 1 3 o r 19. Solubility parameter calculated at 140 K from eq 4. Heat of fusion, taken from ref 13, unless otherwise noted. e Heat of solution, inferred from e q 11 and A P in Table I. f A constant characteristic of the solvent-solute pair. Chosen t o make a H S from eq 6 equal the observed value of A H ? Number density of solute. molecules in saturated liquid krypton at 1 4 0 K. Calculated from eq 5 and 7 and A H S given in this table. Taken from Table I.
I 51
C02 IN LKr v2 BAND
convenient temperature, T = 140 K. The solid lines in Figure 2 correspond to the values of 11140 and A l P in Table I. The solid lines in Figure 3 correspond to the values of AZP in Table I for SOz and CCl,. These lines are visually adjusted representations of the data. Upper limits on nI4 are given for SO2 and CC14. The difference between A l P and AH*may be estimated as follows. The density, p ( T ) , of liquid krypton varies slowly with temperature, T. Set X 2 / = X,(T?, p' = p ( T ? , etc. Since
fl
f\
OF I3c0,
RESOLUTION 0 7 cm-'
I47 2 K
Y
W
0
z a
I O
I
m
a
%
m
I /'J\
/
a
/
650
\
;
1'
I26 9 K
In ( X 2 / X i ) = In ( n / n ? + In ( p ' / p )
Y\',
645
640
(ern-') Figure 4. Spectral profiles of the vP band of '%02in naturally abundanct COPat different temperatures. Absorbance In ( I o / I ) ,where I/I, is the optical transmission through the cell. The solution is saturated with COPat 120.7 and 126.9 K and is unsaturated at 147.2 WAVE NUMBER
K. The small-scale structure in the profiles is instrumental noise. The overlapping absorptions of the "C02 vP band, which peaks near 622 cm-', have been subtracted.
COz absorption feature which was integrated to obtain A is shown in Figure 4 and is typical of the spectroscopic data. The spectral ranges, A?, over which the absorptions were integrated for each solute are given in Table I. We assume that the integrated absorbance, A , is constant at temperatures above the precipitation point, when the measured amount of solute is entirely dissolved. Thus, it has been possible to determine the absolute solubilities of the compounds for which data are given in Figure 2. Only relative solubilities have been determined for SOz and CCl,, for which data are shown in Figure 3. In order to extend the temperature range over which relative solubilities could be obtained for SOz and MoF,, A is measured in two and three spectral ranges, respectively (see Table I), and values of A from different spectral ranges are normalized at those temperatures where spectroscopic data can be conveniently obtained for two spectral ranges. The data in Figures 2 and 3 are tabulated as supplementary material (see paragraph at end of text regarding supplementary material) and can be empirically represented by In (n/nI4,,) = -AHm/RT
(9)
where AZP is a parameter that is approximately equal to AHs as will be shown below, n is the number of solute molecules per cm3 in a saturated solution, and n140is the number of solute molecules in a saturated solution a t a
(10)
it follows that AH"' calculated from eq 9 and n/ni40 will differ slightly from AH*. In principle, if In ( X 2 / X i )plotted against 1/T is a straight line, then In (n/n?will be curved. The curvature would be too slight to be seen in Figures 2 and 3, but a straight line representing the data would be slightly tilted. If, for a given solute, we have data for T' 5 T 5 T", then AH*will be approximately related to AHm in Table I by In
(g) =
The density of liquid krypton varies fromI3 2.04 to about 2.43 g/cm3 between 163 and 119 K; in the case of MoF,, therefore
AH*- A H m = 153 cal/mol or =5% of the value of AH"'. Values of AH*calculated from AH"' in Table I and eq 11 are listed in Table I1 for several solutes; the differences between AHmand AH0 are small, mostly within the experimental uncertainties given in Table I. The integrated absorption cross sections, &, in the indicated spectral ranges, A?, have been calculated from eq 8, the values of A represented by the dashed lines in Figure 2, and the known amounts of MoF,, CO,, HN3, CH20,and CH30H in the cell. Because of the high absorbances for the C 0 2 bands, a was measured for the v2 band of 13C02; the value of Q in Table I was then computed from the measured value of A , the measured solubility of CO,, and the knowni4 natural isotopic abundance (1.11%) of I3C. (13) "Landolt-Bornstein, Zahlenwerte and Functionen", SpringerVerlag, West Berlin, 1961: H. Borchers, H. Hausen, K. H. Hellwege,K. Schafer, and E. Schmidt, I1 Band, 1 Teil, 1971,pp 547,550; 4 Teil, 1967, pp 172,174, 178,200,544,546,574,576,660,717. J. Bartels, H. Borchers, P. Ten Bruggencate,H. Hausen, K. H. Hellwege, K. Schtder, E. Schmidt, I1 Band, 4 Teil, 1961, p 226.
The Journal of Physical Chemistry, Vol. 86, No. 22, 1982 4355
Solubilities of Small Molecules in Liquld Krypton
h
$?
v
=
0 v, v,
o
3800
I
L
3400
~ VI
,
l
l~
,
l
,
l
,
I
,
l
,
2200
2600
3000
1
1600
1400
I200 1000 WAVE NUMBER (cm-')
800
600
Flgure 5. Infrared spectra of CH,OH vapor and CH,OH in solution in liquid krypton (LKr). The molecular density of CH,OH is approximately 1.2 X 10'' molecules/cm3 in both spectra, which is the equivalent of 70 ppm mole ratio in solution. Assignments of the modes of the vapor and pure liquid phases are indicated above and below the spectra, respectively. Labels with (7) are possible assignments for the solution. Other assignments in solution agree with the pure liquid assignments. The CO, bands are due to impurities in the krypton. The spectral resolution of the spectrophotometer is =1.5-3 cm-'.
The integrated cross section, 6.4 X 10-l8 cm2 cm-l-5%, +15%, thus obtained for the v2 band of C 0 2 dissolved in liquid krypton is similar to the value, 7.6 X 10-l8, derived from the absorptivity of C02 vapor.15 The absorption spectrum of methanol dissolved in liquid krypton is shown in Figure 5. Bands appear a t or near the expected positions of the v1 (0-H stretch), v2, v7, and v12 modes of liquid methanol.16J7 Wavenumbers of absorption bands of CH30H vapor are also indicated, and it can be seen that the vl, v7, and v12 bands in Figure 5 are shifted from their gas-phase locations.16
IV. Discussion We now investigate how well eq 5-7 represent our data. Table I1 contains values of the parameters in eq 5-7. It is evident that AHs # AH$ for any of the solutes, so these solutions are not ideal. The parameter ll2 is predicted8 to be In fact, values found for lZlzl typically7-" range as high as 0.1, and 11121 i= 0.18 has been obtained oc~asionally.~ Values of lI2 have been obtained for the solvent-solute systems studied here by setting eq 6 equal to the observed value of AHs and solving for lI2. The values of 2,, required (14) R. C. Weast, Ed., 'Handbook of Chemistry and Physics", 51st ed., Chemical Rubber Publishing Co., Cleveland, OH, 1970. (15) S. S. Penner "Quantitative Molecular Spectroscopy and Gas Emissivities", Addison-Wesley, Reading, MA, 1959. (16) T. Shimanouchi, 'Tables of Molecular Vibrational Frequencies", Consolidated Vol. I, US.Government Printing Office, Washington, DC, 1972, Natl. Stand. Ref. Data Ser. ( U S . , Natl. Bur. Stand.), No. 39. (17) C. J. Pourchert, 'The Aldrich Library of Infrared Spectra", 2nd ed., Aldrich Chemical Co., Milwaukee, WI, 1975.
to fit eq 6 to our AHs range from -0.003 for dissolved COz to 0.427 for dissolved CH30H. As might be expected, the values of 11121 tend to be larger for the more polar solutes. Table I1 also contains vlaues calculated for n140 from eq 5 and 7 and the observed value of AH8, i.e., from A H 8
In x2(140 K) = -- RTo 140R
(12)
n140 = (1.675 X 1022)x2(140 K) cm3 where the molecular density of liquid krypton at 140 K is 1.675 X loz2~ m - ~These . calculated values of n140 are within a factor of 3 of the observed values of n140, except for CH30H. Thus, the theory of Hildebrand, Preston, and Prausnitz,' eq 5-7 and 12, roughly describes the solubility in liquid krypton of MoF6, CC14,COz, and SOz, with 0.003 I )1121 I 0.25. The solubility of CH30H is not even approximately described by eq 12, and in fact the predicted value of nla is lo4times smaller than the observed value. The reason for this failure of the theory is discussed below. The absorption spectrum of dissolved methanol in Figure 5 more closely resembles that of liquid methan01,'~J~ which is known to be associated,6J8 than it does that of gaseous, monomeric CH30H. This spectrum is thus strong evidence that the methanol is in solution in liquid krypton as dimers or higher aggregates. The theoretical treatment
-
(18) J. Errera, R. Gaspart, and H. Sack, J. Chem. Phys., 8,63 (1940). (19) E. W. Washburn, "International Critical Tables of Numerical Data, Physics, Chemistry and Technology", McGraw-Hill, New York, 1926 pp 27-8.
J. Phys. Chem. 1982,86,4356-4360
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herein has thus far not been adequate to deal with associated solutes. The solubility of methanol in liquid krypton can be adequately represented by the theory if it is assumed that the methanol is associated in solution and if one recognizes that the AH2 for methanol in Table I1 is not the appropriate heat of fusion to use in these computations. Since there will be l / k as many (CH30H), aggregates as CH30H molecules in a solution of associated methanol molcules, eq 12 should be rewritten
where k is the average number of CH30H molecules in a dissolved aggregate, provided that the solutions are dilute, which is the case here. (Equation 13 can also be derived from information in ref 6, pp 316 ff and 388-91, for dimers, Le., for k = 2.) The proper heat of fusion to use in eq 13 is the value which takes crystalline methanol to a hypothetical liquid of associated methanol dissolved in liquid krypton. Fairly large aggregates normally exist in liquid methanol,, but small aggregates are more likely to be soluble in cold liquid krypton than large ones. Since the tabulated heat of fusion for methanol is appropriate for melting pure solid CH30H to form pure liquid, and since additional energy will be required to break the large aggregates into smaller ones, it follows that the value of AH$ to be used in eq 13 will be larger than the value given in Table 11. We can use eq 13 to estimate the value of AHof required to take CH30H from its crystalline state to a hypothetical pure liquid consisting of CH30H aggregates having a value
of k identical with that in our solutions. This estimate requires, in principle, that k be known, but the value calculated for AH; does not depend strongly on k . Thus, from xz(140 K) = 1.55 X and AHB= 6343 cal/mol, one obtains AH,' = 3844 cal/mol when k = 2 and 3526 cal/mol when k = 5. Both of these values of AH; are, as expected, considerably larger than the published heat of fusion, 798 cal/mol, given in Table 11. We conclude that the CH30H aggregates dissolved in liquid krypton are considerably smaller than the aggregates in pure liquid CH30H near its freezing point. V. Summary Solubilities of MoF,, COz, CH,O, HN3, and CH30H in liquid krypton have been measured between 118 and 165 K. Relative solubilities of CC14and SOz in liquid krypton have also been measured. The solution theory of Preston and Prausnitz7 roughly describes the solubility of MoF,, CC14, COz, and SOz. Methanol dissolves in liquid krypton as an aggregate, and its solubility can be described by the theory of Preston and Prausnitz7 if it is supposed that the average degree of association of methanol dissolved in liquid krypton is smaller than it is in pure methanol near its freezing point. Integrated cross sections have been measured for a few infrared absorption features of MoF,, COz, CHzO, HN3, and CH30H. Acknowledgment. This work was performed under the auspices of the U.S. Department of Energy. Supplementary Material Available: Table of solubility data (1page). Ordering information is given on any current masthead page.
Polymorphism in the Tetrahalophthalic Anhydrides. 2. The Crystal and Molecular Structures of Ordered Tetraiodophthalic Anhydridet D. S. Sake Gowda and Reuben Rudman' Department of Chemistry. Adelphi University, Garden City, New York 11530 (Received: March 15, 1982; In Final Form: July 28, 1982)
Tetraiodophthalic anhydride (TIPA), like the other tetrahalophthalic anhydrides, undergoes a phase transition near the melting point and forms charge-transfer complexes with polycyclic compounds. The crystal and molecular structures of TIPA have been investigated so that the structural relationship between the two phases and the effects of complex formation on molecular distortion can be determined. TIPA crystallizes as yellow needles, tetragonal system, a = 22.655 (3), c = 9.136 (1)A,2 = 16 molecules/unit cell, D, = 3.69 g/cm3, D , = 3.50 g/cm3, space group I 4 , / a . The structure has been solved by using 1512 unique reflections and refined with least-squares analysis to a final conventional agreement factor of 0.037. TIPA crystallizes in a crystal system different from the chloro and bromo analogues (TCPA and TBPA, respectively) but the molecular shape and distortions from planarity are similar for all three compounds. The structures of these compounds are compared.
Introduction The tetrahalophthalic anhydride (TXPA) compounds, CGX4(CO)20,where X = F (TFPA), C1 (TCPA), Br (TBPA), or I (TIPA), have been studied extensively in recent years for several reasons. Theoretically the molecule should be planar, but the overcrowding of the benzene ring by the halogen atoms results in molecular distortion. The effect of the halogen atom size on this distortion is of 'For part I see ref 4.
0022-3654182/2086-4356$01.25/0
interest. Also, since there is no hydrogen bonding, the crystal packing motif is determined by weak intermolecular forces, which vary with the halogen atom present in the molecule. Finally, the TXPA compounds are known to form charge-transfer (CT) complexes with a number of polycyclic c0mpounds.l It is advisable to determine the structures of the uncomplexed TXPA molecules so that (1) A. K. Wilkerson, J. B. Chodak, and C. E. Strouse, J. Am. Chem. SOC.,97, 3000 (1975); R. D. Srivastava and P. D. Gupta, Spectrochim. Acta, Part A, 24, 373 (1968).
0 1982 American Chemical Society
