1162
F. SCHREINER, G. MCDONALD, AND C. CHERNICK
The Vapor Pressure and Melting Points of Xenon Difluoride and Xenon Tetrafluoride132 by Felix Schreiner, Geraldine N. McDonald, and Cedric L. Chernick Argonne National Laboratory, Argonne, Illinois
60489 (Received August 7 , 1967)
Vapor pressure measurements have been carried out on samples of pure xenon difluoride and xenon tetrafluoride. The preparation and purification of the samples are described as well as the apparatus with which the data for the two reactive fluorine compounds were obtained. For xenon difluoride, the following equation represents the experimental results between 273 and 388°K log P,,
=
--3057'67 - 1.23521 log T + 13.969736 T
The corresponding vapor pressure equation for xenon tetrafluoride valid between 275 and 390.25"K is log P,,
=
3226.21 T
--
- 0.43434 log T
+ 12.301738
In addition, the triple-point temperatures of the two compounds have been determined by a thermal arrest method. For xenon difluoride a triple-point temperature of 402.18'K was found, and for xenon tetrafluoride, 390.25"K. The enthalpies of vaporization at 330°K were derived from the vapor pressure equations and the following values were obtained: AHsub= 55.2 i 0.8 kJ mole-' for XeF2, and AHBub= 60.6 i 1.0 kJ mole-' for XeF4. The standard entropy of xenon tetrafluoride vapor at 298.15'K was calculated from the vapor pressure and heat capacity data and found to be 305.2 J deg-l mole-'. This number is compared with the entropy calculated from molecular data. The entropy of solid xenon difluoride at 330°K was calculated to be 122.8 J deg-1 mole-'.
I n spite of the fact that the properties of chemical compounds of the rare gas xenon have been studied rather intensively in recent years, the literature contains very few data for the vapor pressures of the two lower xenon fluorides. This is, of course, a consequence of the difficulties encountered in attempts to carry out precise measurements a t elevated temperatures with these reactive substances. The available data were obtained around room temperature from direct measurements of the pressure with h!Ionel Bourdon gauges. For xenon difluoride, a pressure of 3.8 mm at 25" was observed by Agron, et aLj3 and for xenon tetrafluoride, Claassen4 found approximately 3 mm a t 20". The accuracy of these values is approximately h 0 . 5 mm, based on the reading accuracy of the Bourdon gauges. Both the difluoride and the tetrafluoride have also been the subject of a study of the variation of the vapor pressure with temperature.5 These measurements were carried out spectroscopically, the chief aim being the evaluation of the slope of the vapor pressure curve in order to estimate the enthalpies of sublimation. I n the case of xenon difluoridebthe intensity of the characteristic absorption a t 1750 A was followed between - 15 and 22". For the tetrafluoride, similar absorption measurements were made at 1330, 1586, and 2010 8 a t temperatures between -3 and +22". The measured variation of the absorption with temperature was used
+
The Journal of Physical Chemistry
to calculate the enthalpies of sublimation from the Clausius-Clapeyron equation, and the values 51.5 and 64.0 kJ mole-' were obtained for the difluoride and tetrafluoride, respectively. The scarcity of the available melting point data for the two xenon fluorides equals that of the vapor pressure data. For xenon difluoride, two values have been , ~ more recently reported, 140" by Agron, et ~ l . and 130" by Gr6z, et aL6 The only piece of information on record for xenon tetrafluoride is an approximate value of 114"for the melting point, observed by Chernick.' I n view of the exiguity of the existing information, the necessity was felt for systematic vapor pressure measurements covering a more extensive range of temperatures and pressures. The results of such studies for pressures up to 1 atm are reported in the present (1) Based on work performed under the auspices of the U. S. Atomic Energy Commission. (2) Presented at the 153rd National Meeting of the American Chemical Society, Miami Beach, Fla., April 9-14, 1967. (3) P. A. Agron, G. M. Begun, H. A. Levy, A. A. Mason, C. F. Jones, and D. F. Smith, Science, 139, 842 (1963). (4) H. H. Claassen, H. Selig, and J. G. Malm, J . Am. Chem. SOC.,84, 3593 (1962). (5) J. Jortner, E. G. Wilson, and 8. A. Rice, ibid., 85, 814 (1963). (6) P. Gr6z, I . Kiss, A. RBvBsz, and T. Sipos, J. Inorg. Nucl. Chem., 28, 909 (1966). (7) C. L. Chernick, "Noble Gas Compounds," University of Chicago Press, Chicago, Ill., 1963, p 35.
VAPORPRESSURES AND MELTINGPOINTS OF XeF,
AND
1163
XeF4
paper, together with the results of triple-point determinations of the two lower xenon fluorides.
Experimental Procedure Preparation of Xenon Dijiuoride. Xenon difluoride was prepared by heating a mixture of 0.79 g-atom of high purity xenon and 0.08 mole of fluorine in a 574-cm3 nickel reaction can at 300" for 24 hr. According to the equilibrium constants for the homogeneous vapor phase equilibrium of the xenon-fluorine system,* the formation of only small amounts of xenon tetrafluoride and negligible amounts of the hexafluoride is expected under these conditions. Purification from XeF4 and XeF6 could be achieved by stepwise removal of the vapor above the product. After subliming off a total of 0.2% of the solid, the infrared spectrum obtained showed only bands attributable to difluoride, and the product was judged sufficiently pure for the vapor pressure and melting point determinations. Preparation of Xenon Tetra$uoride. From an inspection of the vapor phase equilibrium data, it is clear that the tetrafluoride formed in a homogeneous gas reaction will always be accompanied by appreciable quantities of difluoride, or of hexafluoride, or of both. I n order to circumvent this difficulty, the sample was prepared in a system where the solid product could condense out as the reaction proceeded. In principle, this method was suggested by Weinstock, et ~ l . and , ~ has been adapted here for application at lower temperatures. A 1600-cm3 nickel reaction can was filled with 0.68 mole of fluorine and 0.30 g-atom of xenon. The upper part of the can was fitted with a sleeve heater surrounded with asbestos tape for thermal insulation. For temperature measurement and control, a chromel-alumel thermocouple was inserted between heater and can. The temperature of the top part of the can was then raised to 400" and kept there for 48 hr in order to let the mixture react. At the same time, the lower part of the can was immersed in a water bath to a depth of about 5 cm. Throughout the reaction time, the temperature of the water bath was kept below 50'. As soon as the pressure of a reaction product exceeds the vapor pressure at the low temperature a condensed phase will be formed in this system. Barring the existence of solid mixtures, only one solid compound can be present a t the end of the reaction, and it has to consist of either the pure difluoride or tetrafluoride, depending on the initial proportion of the reactants. I n the present case, an atomic ratio F/Xe of 4.5 to 1 was chosen in order to form the tetrafluoride. At the end of the experiment, the reaction can was cooled to 77'11, and the unreacted fluorine was removed. It turned out that about 20% of the initially supplied fluorine had not been consumed. After warming to 195°K it was found that practically all of the xenon had reacted. An infrared spectrum of the vapor above the product
showed traces of difluoride and hexafluoride to be present. By stepwise removal of the vapor, the hexafluoride could be removed from the sample before 1%of the product had been sublimed. The last traces of difluoride, however, disappeared only after about 20% of the sample had been transferred to a second storage can. The tenacity with which the difluoride was retained in the tetrafluoride can be accounted for in two ways: either the evaporation of the difluoride was strongly inhibited by occlusion in tetrafluoride crystals or the difluoride was present as the known 1 :1 compound with a much lower partial pressure.1° Vapor Pressure Apparatus. The measurements of the vapor pressure were performed on samples ranging in mass from 2 to 3.5 g, contained in the bottom part of a vertically mounted closed-end nickel tube of 300 mm Iength. The tube was attached to a diaphragm-type pressure transducer which has been described in detail by Sheft.ll Those parts of the system with which the sample came into contact during the measurements, nix., sample tube, pressure transducer, and connecting valve, were made of nickel or Monel and had been passivated by previous treatment with C1F3 and XeF6. The upper part of the apparatus was contained in a heatable box which was kept at a temperature exceeding the highest temperature reached during each series of runs by about 20". Pressures were determined by operating the diaphragm transducer as a null indicator. A nitrogen pressure equal to the vapor pressure was applied to the reference side of the transducer and measured with a fused quartz Bourdon gauge that had been calibrated against a mercury manometer to give pressures in millimeters of mercury at 0" and standard gravity. The experimental error of the pressure mea= surement was determined chiefly by the instability of the zero adjustment of the transducer, which produced anuncertainty of h0.2 mm. The lower half of the sample tube was immersed in a bath whose temperature was controlled to several hundredths of a degree. Temperatures were measured with an NBS calibrated platinum resistance thermometer, the resistance of which was determined with a Mueller resistance bridge. Triple-Point Temperatures. The triple-point temperatures of xenon difluoride and of xenon tetrafluoride were determined by a thermal arrest method on samples of 3.75 and 2.06 g, respectively. The measurements were carried out in nickel tubes of 57 mm length, 9.5 mm o.d., and 0.38 mm wall thickness. The tempera(8) B. Weinstock, E. E. Weaver, and C. P. Knop, Inorg. Chem., 5 , 2189 (1966). (9) B.Weinstock and E. E. Weaver, presented a t the 149th National Meeting of the American Chemical Society, Detroit, Mich., April 1965. (10) J. H.Burns, R. D. Ellison, and H. A. Levy, J. Phys. Chem., 67, 1569 (1963). (11) I. Sheft, Rev. Sci. Instr., 37, 767 (1966).
Volume 78, Number 4
April 1968
1164 ture was measured with a copper-constantan thermocouple which had been calibrated against an NBS calibrated platinum resistance thermometer. The thermocouple junction was located near the bottom of the sample tube inside a center thermocouple well of 1.6 mm 0.d. Cooling and warming curves were recorded on a strip chart recorder connected to the potentiometer with which the emf of the thermocouple was measured. The samples were placed in a furnace and observations were made at cooling and warming rates of 0.5" min-1. The thermal arrest times were between 15 and 20 min, and the temperature variation during the first half of these periods was at most 0.02" for both samples. I n two independent determinations the triple-point temperature of xenon difluoride was found to be 402.18 f O.05"K. Similarly, the triple-point temperature of xenon tetrafluoride was observed to be 390.25 i= 0.05"K. The value for XeF2 agrees within the experimental error with the temperature reported by Gr6z, et aL6 For XeF4 the new value is 3.1" higher than the one observed by Chernick.'
F. SCHREINER, G. MCDONALD, AND C. CHERNICH logP,,
=
The Journal of Physical Chemistrg
T
1.23521 log T
+ 13.969736
(1) I n the derivation of this equation, empirical weights were applied to take account of the fact that at higher temperatures the pressure deviations increase because of the growing importance of the temperature fluctuations of the bath. Equation 1 represents the experimental data in the temperature range from 273 to 388°K within the limits of the experimental error. Figure 1 shows a plot of the deviations of the observed pressures from the ones calculated from the equation. It is not possible to represent the data by a two-constant equation of the type log P,, = -A / T B because the variation of the enthalpy of sublimation over the temperature range of the measurements produces a marked curvature. The number of data points available for xenon tetra-
+
L
+LOO/
Results and Discussion Vapor Pressures. The vapor pressure measurements on the two lower fluorides of xenon were carried out over a temperature interval of more than 100" and cover a pressure change of three orders of magnitude. All values were corrected for the pressure head of the column of vapor in the sample tube and the column of nitrogen in the reference line in order to obtain the pressures a t the interface of the condensed and vapor phases. This correction amounts to at most 0.14 mm a t 1 atm and is proportionately smaller a t lower pressures. No correction was applied for dissociation of the compounds in the vapor phase. The maximum error arising from neglecting this correction is 0.6 mm for XeF4 at the highest temperature. Since the reaction rates are known to be slow, however, the extent to which the equilibrium was established during the time of the vapor pressure measurements is uncertain. I n addition to the vapor pressure data measured with the nickel apparatus described above, two series of earlier measurements had been carried out in a system employing a stainless steel diaphragm transducer. This system did not permit the extension of the measurements beyond 314°K and placed a severe restriction on the measurable pressure range. The data obtained with the stainless steel transducer, however, are in quite good agreement with the more recent measurements and have therefore been used in deriving the vapor pressure equations. It is worth noting that the agreement between the two sets of measurements also attests to the purity of the substances, since the samples originated from different preparations. For xenon difluoride a total number of 87 data points were available. These were fitted by a least-squares method to yield the equation
--3057.67 -
I
I
$0.50 0
1
r
L I
I
] I
I
1
E
n'
0 0 '
oaoo
a
I
-0.50
I
i
270
290
-1.00'
I
I
I
310
I
I
I
330
T,O K
I
I
i " f . 1
I
I
I
350
370
390
Figure 2. Plot of the deviations of the observed vapor pressures of xenon tetrafluoride from the pressures calculated from eq 2. Filled-in circles refer to data obtained with a stainless steel apparatus.
VAPOR PRESSURES AND R'IELTING POIXTS OF
fluoride was 58. They are represented within the limits of error by the equation logP,,
=
3226.21 T 0.43434 log T
1165
XeFz AND XeF4
-~
+ 12.301738
(2)
This equation is valid for temperatures ranging from 275°K to the triple point for fusion, 390.25"K. Figure 2 shows a plot of the deviations of the observed and calculated pressures. The scatter of the points on the two deviation plots (Figures 1 and 2) is around several tenths of a millimeter at low pressures where the error of the pressure measurement determines the experimental error, and increases to about 1 mm at the highest temperatures, where the fluctuations of the bath temperature become more important. A single point was measured for liquid xenon tetrafluoride at 390.51°K, and the vapor pressure was found to be 817.97 mm. A comparison of the vapor pressures of the two compounds shows that throughout the entire temperature range, xenon difluoride is the more volatile substance. At 298.15"K the pressure of the difluoride is 4.55 mm, whereas the tetrafluoride has a vapor pressure of only 2.55 mm. Enthalpies qf Sublimation, The enthalpies of sublimation of XeF2 and XeF4 were calculated from the Clausius-Clapeyron equation for a temperature of 330"Ii, and the following two values were obtained: AHsub = 55.2 3t 0.8 kJ mole-' for XeF2, and AHsut, = 60.6 f 1.0 k J mole-' for XeF4. The limits of error are based on a careful consideration of the maximum uncertainty of the pressures as reflected by the deviations of the measured points from the vapor pressure curve. I n deriving the enthalpies of sublimation, the ClausiusClapeyron equation was considered adequate because of the low vapor pressure values at 330°K (39.05 mm for XeFzand 27.01 mm for XeF4). I n comparing the sublimation enthalpies with the values published previously by Jortner, et aL16attention has to be paid to the fact that these latter data refer to a n average temperature of 277°K. At this temperature, the enthalpy of sublimation of xenon difluoride derived from eq 1 has the value 55.7 kJ mole-'. Jortner, et ai.,5 found 51.5 f 0.8 kJ mole-' from their spectroscopic studies. The difference between the two numbers exceeds the limits of error considerably. It has been pointed out by Jortner, et al.,5 that the low volatility of both xenon fluorides implies a high lattice energy of the solid which cannot be accounted for by van der Waals forces alone. Consequently, the bonds in these compounds must have an appreciable polar character. For this reason, Jortner, et ak., proposed a model for the bonding in the xenon fluorides which involves a substantial migration of charges to the ligand atoms. It was estimated that the charge carried by a
fluorine atom in XeFz amounts to 0.5 elementary charge. A similar estimate for XeF4 resulted in the transfer of 0.42 elementary charge. A calculation of the enthalpy of sublimation based on this model yielded a value of 55.6 kJ mole-' for crystalline xenon difluoride, very close to the value obtained from eq 1. I n the case of xenon tetrafluoride, the enthalpy of sublimation at 277°K was not obtained directly from the slope of eq 2 a t that temperature. Since 277°K is rather close to the low end of the range of temperatures for which the vapor pressure equation is valid, it was preferable to correct the more reliable value for AHsub at 330°K to the lower temperature. The enthalpy change between 330 and 277°K was calculated from the difference of the heat capacities of gaseous and solid XeF4. The vapor heat capacity at 298.15"K was obtained from molecular data and has the value 89.8 J deg-1 mole-'. The heat capacity of the solid at the same temperature is known from the measurements of Johnston, et ai.,12 to be 118.5 J deg-l mole-'. With these numbers the enthalpy change between 330 and 277°K is estimated to be 1.5 k J mole-', and one obtains a value of 62.1 & 1.0 kJ mole-' for the enthalpy of sublimation at 277°K. This is only slightly lower than the value of Jortner, et aL5 Entropies. The standard entropy of xenon difluoride vapor at 330°K calculated from molecular data is 265.2 f 0.5 J deg-' mole-'. This number is based on the frequency assignment of Agron, et aL13 and on the assumption that the Xe-F bond distance is equal to the Xe-F distance in crystalline XeF2.13 Since no lowtemperature heat capacity data are available, the calculated vapor entropy cannot be compared with a value obtained from calorimetric measurements. However, by adding 24.7 J deg-' mole-' for expansion of the vapor, and subtracting the entropy of sublimation, 167.1 J deg-' mole-', a value of 122.8 3t 3.0 J deg-1 mole-' can be given for the entropy of solid XeFz at 330°K. The comparison of the calculated vapor entropy with the calorimetric value is of particular interest in the case of xenon tetrafluoride because of the unusual planar configuration of the molecule. The pertinent data for such a comparison are listed in Table I. The entropy of solid xenon tetrafluoride at 298.15"K was taken from the heat capacity work of Johnston, et a1.l2 The entropy of sublimation was obtained from a value of the enthalpy of sublimation adjusted to 298.15' K as described before. No correction for nonideality of the vapor was applied because the vapor pressure of the compound is only 2.55 mm. The calculated entropy value is based on the following set of molecular data. The bond distance of 1.94 A was taken from the (12) W. V. Johnston, D. Philipovich, and D. E. Sheehan, "Noble Gas Compounds," University of Chicago Press, Chicago, Ill., 1963, p 139. (13) H. A. Levy and P. A. Agron, J. Am. Chem. SOC.,85,241 (1963).
Volume 72, Number 4
April 1968
JOHN T.SEARS AND JAI\fES w. BUTHERLAND
1166 electron diffraction work of Bohn, et ~ 1 . ~ For 4 the frequencies of the seven normal modes of vibration, the assignment of Claassen, et uZ.,'5 was chosen, except for the value of v4, which is infrared and Raman inactive and which was taken from a publication of Yeranos,16 and the value of v7, for which the assignment of Weinstock, et U Z . , ~ was adopted. The individual frequencies used in the calculation are: VI = 550 cm-', v2 = 291 cm-', v3 = 235 cm-', v4 = 231 cm-', 215 = 519 cm-', Most of these V6 = 586 cm-l, and ~7 = 250 cm-'. frequencies are nondegenerate, and only V6 and v7 are doubly degenerate. The assignment of the lowest frequency, v7, might be expected to be as low as 167
Table I: The Standard Entropy of Xenon Tet'rafluoride Vapor at 298.15"K (Molecular Weight: 207.29)
Entropy of solid XeF4 Entropy of sublimation,
146.4 rfi 4 . 2 J deg-1 mole-' 206.2 rfi 5 . 0 J deg-I mole-'
61480/298.15
Entropy of compression, -R In 760/2.55 Standard entropy of gaseous XeF4 Standard entropy of gaseous XeF4, calculated from molecular data
cm-' on the basis of force field consideration^,'^ but the value chosen here is more consistent with the calorimetric entropy. As can be seen from the table, the calorimetric entropy differs from the calculated one by 11.4 J deg-' mole-l. This difference is barely within the experimental error, but since the limits of error have been assigned rather carefully, the discrepancy appears to be uncomfortably high. I n order to obtain complete agreement between the two numbers, for instance, the frequency of v7 would have to be increased to 300 cm-l, contrary to all expectation. It seems, therefore, that consideration should be given to the acquisition of more reliable thermodynamic and molecular data of xenon tetrafluoride in order to improve the internal consistency of all available data.
Acknowledgments. The authors wish to express their gratitude to Drs. D. W. Osborne and H. H. Claassen for fruitful discussions and helpful suggestions in the preparation of this paper.
-47.4 Z!Z 0.7 J deg-1 mole-' 305.2 i. 10 J deg-1 mole-' 316.6 Z!Z 1 . 5 J deg-l mole-'
(14) R. K. Bohn. K. Katada. J. V. Martinez. and 5. H. Bauer. "Noble Gas Compounds," University of Chicago Press, Chicago; Ill., 1963,p 238. (15) H. H. Claassen, C. L. Chernick, and J. G. Malm, J. Am. Chem. SOC.,85, 1927 (1963). (16) W.A. Yeranos, Mol. Phys., 9, 449 (1965). (17)H.H.Claassen, private communication.
Radiolytic Formation and Decomposition of Ozone' by John T. Sears and James W. Sutherland Brookhaven National Laboratory, Upton, N e w York
11978
(Received August 7, 1967)
The formation of ozone from gaseous oxygen upon irradiation with Corn y rays was investigated for various dose rates, I , ranging from 7 x loa t o 8 X 106 rads/hr. Yield-dose curves at ambient temperature ( 5 10") ~ and 665 torr were measured until a steady-state ozone concentration was reached. G(O1) at low doses was found to be independent of I. Ozone concentrations of 2000 ppm at ambient temperatures and 7000 ppm at 195'K were detected. In the gas phase at 77°K) a differential G(O1) of 10.5 =t0.5 was found. The concentration at steady state was determined by a separate back reaction(s) which destroyed ozone and by the forward rate of formation of ozone from its precursors. Different dose-rate effects were found under various conditions. The relation, [08]38 0: 1°,45*0.08, was determined for the steady-state yield in all-glass vessels, while [O,lB. 0: P.g1*0.04 held when the irradiation vessel was fitted with a fluorocarbon-lubricated stopcock. Both nonradiolytic and radiolytic decomposition of ozone were also investigated.
Introduction The formation of Ozone from oxygen, one of the oldest radiation chemistry Systems, was first investigated in 1911 by Lind.2 Recent worker^^,^ found that ozone was formed more efficiently at low temperatures than at room temperature and that only small concentrations T h e Journal of Physical Chemistry
of a few ppm were attained at room temperature. Earlier work5~6showed that decomposition of ozone proceeds at an extremely rapid rate under irradiation. (1) This work was performed under the auspices of the U. 8. Atomic E~~~~~commission. (2) s. c. Lind, Amer. Chem. J., 47, 397 (1911).
