952
MAXT. ROGERS AND EMERSON E. GARVER
Vol. 62
VISCOSITIES AND SURFACE TENSIONS OF SOME LIQUID HALOGEN FLUORIDES'*' BY MAXT. ROGERS AND EMERSON E. GARVER Department of Chemistry, Michigan State University, East Lansing, Michigan Received March $0, 1968
Values of the surface tension and viscosity of iodine pentafluoride, bromine pentafluoride and bromine trifluoride have been measured in the liquid phase over a range of temperature with a precision of i 2 % and the data have been fitted to empirical equations. By application of the Eyring theory of viscosity the heats and free energies of activation for viscous flow have been computed and are compared with the energies of vaporization. These results, along with the Trouton constants and the dielectric constants, indicate that iodine pentafluoride and bromine trifluoride are associated in the liquid phase, whereas chlorine trifluoride and bromine pentafluoride probably are normal.
Introduction Although many physical properties of the halogen fluorides are now well-known' the only values of surface tensions or viscosity coefficients reported are those for chlorine trifluoride.3 Since these properties are of both theoretical and practical interest for this series of compounds we have made measurements of the surface tensions and viscosities of iodine pentafluoride, bromine pentafluoride and bromine trifluoride over a range of temperature.
Experimental
with carefully purified water, benzene and carbon tetrachloride as standards. The calibration was made with various levels of liquid in the container since this level varied in different experiments. Viscosity coefficients were computed from the observed times of flow t and densities d by use of the simplified equation 11 Adt Bd/t which provides a kinetic-energy correction. Capillary-rise was measured directly with a cathetometer. The radius of each capillary was determined by measuring the length of the thread of a weighed globule of mercury a t various positions in the capillary using the technique of Harkms .e Experimental Data.-Representative values of the viscosities (in centipoise) for the various liquids measured are shown in Table I a t several temperatures over the range covered in each case.
-
Materials.-The bromine fluorides from Harshaw Chemical Co. and the iodine pentafluoride from General Chemical Co. were purified by distillation in a fluorothene column with Monel packing. The amount of impurity was estimated from cooling curves4 to be about 0.05 mole yo for iodine TABLE I pentafluoride and about 0.40 mole % for bromine pentafluoride. The impurity content of the bromine trifluoride VISCOSITY COEFFICIENTS FOR THE HALOGEN FLUORIDES AT was not estimated by this method but is probably less than VARIOUS TEMPERATURES 1 mole 70. The liquids were in every case distilled directly Iodine pentafluoride Bromine pentafluoride Bromine trifluoride into the apparatus thus providing one additional stage of t, 1, 'I? t. "C. 0.p. OC. 0.p. OC. C.P. purification. Apparatus.-Both surface tension and viscosity were 18.9 2.490 2.3 0.824 13.2 3.036 measured in a single apparatus patterned after that of 25.5 2.191 12.3 0.683 20.5 2.553 Doescher and Elvrum.6 A precision-bore Pyrex ca illary, 32.9 1.812 24.2 0.620 25.0 2.219 about 0.75 mm. i.d., was attached to a reservoir o r a b o u t 3 4 . capacity. The reservoir and capillary were sealed 37.7 1.726 28.9 0.590 39.6 1.775 to a large female standard-taper joint. The matching male 40.0 1.686 joint was sealed off to provide a container for the liquid. When the pair was assembled the bottom of the capillary Representative values of the surface tensions of the came near the bottom of the container and the outlet a t the liquids studied are shown in Table I1 a t various temperatures. top was connected to a vacuum line. A thermocouple well was also provided. For measurements on bromine triTABLE I1 fluoride a Vycor apparatus of similar design was constructed. SURFACE TENSION VALUES FOR THE HALOGEN FLUORIDES Fluorothene was tried initially but the contact angle for the Iodine pentafluoride Bromine pentafluoride Bromine trifluoride li uids on fluorothene was found to be fairly large and very d k c u l t to meaaure precisely. Temperature control was t, dyT& 1, dy?es/ t. dykd provided by circulating water from a large constant-temOC. om. OC. cm. om. OC. erature bath around the above apparatus contained in a 12.0 37.1 9 . 2 24.3 18.4 30.8 6ewar flask. The temperature of the bath could be varied 18.9 36.4 14.7 23.5 29.7 25.2 between 0 and 50' and held to &0.05". 27.1 35.6 38.2 28.2 27.0 22.4 The gas-handling system was entirely of Monel and copper except for the vessel described above. The Pyrex (or Vycor) 32.6 21.6 45.0 33.8 apparatus was pre-fluorinated with chlorine trifluoride then the substance to be measured was introduced by distillation. Discussion The time required for the liquid to flow between two marks above and below the reservoir was measured with a stopViscosity coefficients for most liquids obey the watch. The apparatus was calibrated a t each temperature empirical equation (1) fairly well over reasonable r),
(1) Physical Properties of the halogen fluorides. XI. For previous article of this series see M. T. Rogera. J. L. Speirs and M. B. Panish, THIEJOURNAL, 61, 366 (1957). (2) Presented at the 131st National Meeting of the American Chemical Society, Miami, Florida, April 8, 1957. (3) A. A. Banks, A. Davies and A. J. Rudge, J . C h e n . SOC.,732 (1953); since our work was completed a report on the viscosity of iodine pentafluoride has appeared, G. Hetherington and P. L. Robinson, ibid.. 3681 (1956). (4) M. T. Rogers, J. L. Speirs, H. B. Thompson and M. B. Panish, J . Am. Chem. SOC.,18,4843 (1954). (5) R. N. Doescher and G. W. Elverurn, J . Chem. Phus., 40, 834 (1952).
vt
9 =
AeB/T
(1)
temperature ranges. We have fit our data to this equation by the method of least squares and list the values of A and B in Table I11 found by use of all our results (only a small portion of which is shown in Table I). Our values for the viscosity coefficients of iodine pentafluoride agree with those (6) A. Weissberger, Ed., "Techniques of Organic Chemistry," Vol. I, Chap. IX, Second Edition, Interscience Publiehers, New York. N. Y.,1949.
.
VISCOSITIES AND SURFACE TENSION OF LIQUID HALOGEN FLUORIDES
August, 1958
953
TABLE I11 VISCOSITY CONSTANTS FOR A d
Bd
THE
HALOGEN FLUORIDES
CIFab
BrFs
IF,
BrFs
1.48 x 10-4 987
1.11 x 10-4 1195 6.69 2.38 2.76 -1.31 2.43 2.80
4.7 x 10-6 1834 9.28 3.62 3.51 0.39 2.64 2.50
3.41 X 10-6 1944 9.67 4.03 3.35 1.73 2.88 2.50
Quantity"
(3.01 A E , ~ ~(kcal./mole) ' 1.95 AH* (kcal./mole) AF* (kcal./molc) 2.32 -1.22 AS* (e.&) 2.59 A E ~ A F * 3.07 A E,.,/AH * The molar energy of va oriaation of the liquid is AEvap.and the free energy, heat and entropy of activation for viscous flow are Ab'*, AH* and AS P, respectively. * Computed from the data of ref. 3. These are mean values over a temperature range up to the boiling point since the heats of vaporization all are derived from vapor-pressure data. See equation 1. Q
TABLBIV SURFACE TENSION CONSTANTS, MOLAR ENTROPIES OF VAPORIZATION, DIELECTRIC AND CORRELATION FACTORS FOR T H HALOQEN ~ FLUORIDES CONSTANTS Substance
D
Cb
AHvnp/T
e
8
23.1' 4.286d 1.31* Chlorine trifluoride" 26.7 0.160 23.3 7.90 1.06 Bromine pentafluoride 25.4 .113 33.0 .131 26.4 36.2 1.50 Iodine pentafluoride 25.6 ... ... Bromine trifluoride 39.8 .lo4 * The symbols C, D are the constants of equati6n 5. The heats of va orization 0 Computed from the data of ref. 3. of chlorine trifluoride and bromine trifluoride are calculated values a t the normal boiling points; J. N. Grisard, $A. Bernhardt and G. D. Oliver, J . Am. Chem. Sqc., 73, 5725 (!951); G. D. Oliver and J. W. Grisard, ibid., 74, 2705 (1952). The values for iodine pentafluoride and bromine pentafluoride are mean values for the temperature ranges 30-60" and 25-40', respective1 M. T. Rogers, J. L. Speirs, H. B. Thompson and M. B. Panish, J . Am. Chem. Soc., 76, 4843 (1954); M.T. Rogers an8$. L. Speirs, THISJOURNAL 60, 1462 (1956). Only in the case of chlorine trifluoride was the exact Clapeyron equation, with correction for gas imperfection, used. The vapors of iodine pentafluoride and bromine pentafluoride do not RI. a pear to be very imperfect at these temperatures and pressures; no data are available for bromine trifluoride vapor. $ Rogers, H. B. Thompson and J. L. Speirs, J . Am. Chem. SOC.,76, 4841 (1954)and later pa ers of this series. e M. T. Rogers, unpublished calculations; the g factors reported previously (ref. 9)are on the basis of earfikr theories and are believed to be less significant than the present values.
whereas for associated liquids it. is smaller. It is seen from Table I11 that the ratio AEvap/AH* is near three for chlorine trifluoride and bromine pentafluoride but is much smaller for iodine pentafluoride and particularly for bromine triflucride. These ratios along with the positive entropies of hN =: -eAF*/RT (2) activation for viscous flow obtained for the latter V compounds indicate that they are a,ssociated liquids. where AF* is the free energy of activation for The surface tension data were fitted to equation viscous flow, B is the molar volume and the other ( 5 ) , where y, dynes/cm., is the surface tension a t constants have the usual significance. Since
of Hetherington and Robinsona within the limits of experimental error. Eyring and co-workers' have applied the theory of absolute reaction rates t o viscous flow and obtain a relation
,
AF* = AH*
- TAS*
?=C-Dt
this equation may also be written (3)
and, since V and AS* are nearly constant over the small temperature range (2030") studied, this equation may be equated t o equation 1with B = AH*/R (4) where AH* is the activation energy for viscous flow. Vahes of AF*, AH* and of AS* obtained from equations 1 t o 4 are shown in Table 111 for each substance. The molar energies of vaporization (avap) derived from the known heats of vaporization' also are given for each liquid. Eyring, et u Z . , ~ have shown that for nearly all liquids AE,,,./AF* is about 2.45; this is seen from Table I11 t o be true also for the halogen fluorides. Eyring? has further shown that for non-associated liquids AEva,/AH* usually lies between 3 and 4 (7) 8. Glasstone, I(. Laidler and H. Eyring, "The Theory of Rate Processei," McGraw-Hill Book Co.. New York, N. Y., 1941.
(5)
t o and C,D are constants found by application of the method of least squares to all our measured values for a given compound. Values of C and D for the halogen fluorides are shown in Table IV. The relatively high surface tensions indicate strong intermolecular forces in bromine trifluoride and iodine pentafluoride. Other conventional criteria of association in the liquid state are the Trouton constant and the dielectric constant of the liquid. The molar entropy of vaporization of a liquid a t the normal boiling point (Trouton constant) has a value in the range 20-23 for most liquids but is somewhat larger for associated liquids (26.0 for water). Some values of A H v a p / T for the halogen fluorides are shown in Table IV and it is seen that both iodine pentafluoride and bromine trifluoride would be classed as associated liquids on the basis of these data. The measured dielectric constants of the liquids
954
IANM. CROLLAND ROBERT L. SCOTT
and dipole moments of the gaseous molecules8 have been combined with density data t o calculate the correlation factors g for these liquids by use of the theory of Harris and Alder.g The factor g would be unity in the absence of specific orienting forces and, for most non-hydrogen-bonding liquids, is found to lie in the range 0.5 to 1.4. Hydrogenbonded liquids may have much larger values of g (H202, 2.0; HzO, 2.5). The value of g (1.50) found for iodine pentafluoride indicates that there are strong intermolecular forces tending to line up the dipoles in the liquid. The dielectric constant of bromine trifluoride unfortunately is not known. The value of 1.06 found for g for bromine pentafluoride shows that it is not an associated liquid while the value g = 1.30 for chlorine trifluoride is intermediate but low enough t o be accounted for by assuming a non-associated liquid with molecules departing considerably from spherical shape. I n a recent series of crystal structure analyses of complex antimony fluorides evidence for fluorine bridges has appeared.1° Thus in KSbF4 each antimony atom is surrounded by five fluorine (8) M. T. Rogers, H. B. Thompson and J. L. Speirs, J. A m . Chem. Soc., 76, 4841 (1954); M. T. Rogers, R. D. Prnett and J. L. Speirs, ibid., 71, 5280 (1955); M. T. Rogers, R. D. Pruett, H. B. Thompson and J. L. Speirs, ibid., 78, 44 (1956). (9) F. E. Harris and B. J. Alder, J. Chem. Phyr., 21, 1031 (1953).
(IO) A. Bystriim, Arkiu Kemi, 8 , 17, 373, 461 (1951).
Vol. 62
atoms, three a t 2.02 k. and two at longer distances (the mean is 2.35 k.). Each of the latter fluorines is also close t o an antimony atom of another SbFs grouping, the average distance being 2.35 8. Four of these SbFs groupings are bound into an (Sb4Fls)-4 complex in the crystal by these secondary "fluorine bonds." It seems quite possible that similar weak intermolecular bonds are responsible for the strong intermolecular forces in liquid iodine pentafluoride and bromine trifluoride, and for the fact that these liquids show anomalies in physical properties similar to those found for hydrogenbonded liquids. Since the F- ion would play a role here analogous to that of the proton in a hydrogen bond, an additional stable orbital on the central atom would be required. It is indeed observed that the order of stability of the complexes KBrF4 > KIFe > KBrFa
> KCIFl (unknown)
is the same as the apparent order of strength of intermolecular bonding in the liquids BrFs
> IF6 > BrF6 > CIFs
Iodine heptafluoride, which has no extra stable orbital, forms no complex and has none of the properties of an associated liquid. Acknowledgment.-This work was supported by the Atomic Energy Commission through Contract AT-( 11-1)-151.
FLUOROCARBON SOLUTIONS A T LOW TEMPERATURES. 111. PHASE EQUILIBRIA AND VOLUME CHANGES I N THE CH,-CF, SYSTEM BY IANM. CROLLAND ROBERT L. SCOTT Department of Chemistry, University of California, Los Angeles, California Received March S, I968
Liquid-liquid and liquid-solid phase equilibria in the system methane-perfluoromethane have been studied and the phase diagram constructed. The critical solution temperature is 94.5"K., at a critical mole fraction of CF, of approximately 0.43; these values are consistent with the free energies of mixing previously re orted by Thorp and Scott. The eutectic temperature is 84.5"K. a t x = 0.88. The volume change on mixing CH, and 8F4has been measured dilatometrically at 107°K. At x = '/z the volume change is 0.88 ~ mper. mole ~ or 2.1%. This is smaller than the values observed for fluorocarbon-hydrocarbon mixtures at room temperature (ca. 5 cm.3 per mole); when allowance is made for the smaller excess free energy and the small coefficient of thermal expansion of CF,, the whole series fits together consistently.
Introduction I n recent years, much attention has been focussed upon the unusual solvent properties of fluorocarbons and related substances. The low solvent power of fluorocarbons for many organic liquids and solid solutes can be attributed t o their very low solubility parameters (5.7-6.0 cal.1/ecm.-8/p for fluorocarbons liquid a t room temperature) ; many of these solutions are in good agreement with the predictions of Hildebrand's solubility parameter theory. However, solutions of fluorocarbons with aliphatic hydrocarbons are in marked disagreement with expectation. The difference in solubility parameters is such that these solutions should show only small positive deviations from ideal behavior, yet partial miscibility is observed in nearly all such
systems. For example, for the heptane-perfluoroheptane system, one calculates that the critical solution temperature should be about 95°K; it is observed t o be 323'K3 Similar results are observed for all the other pairs studied. Various special explanations have been suggested by different workers, but none of.these-either alone or in combination with others-is entirely satisfactory. In earlier work from this Laboratory4 the prototype system methane-perfluoromethane was investigated in the region of 110°K. The two liquids were completely miscible a t this temperature, but very large positive deviations from ideality were observed. At a mole fraction of one half, the excess free energy F E was found to be 86 cal. per mole; the value estimated from the solubility parameter difference was only 7 cal. per mole.
(1) J. H. Hildebrand and R. L. Scott, "Solubility of Nonelectrolytes," 3rd Edition, Reinhold Publ. Corp., New York, N. Y., 1950. (2) R. L. Scott, THIEJOURNAL, 62, 136 (1958).
Soc., 72, 4348 (1950).
(3) J. H. Hildebrand, B. B. Fisher and H. A. Benesi, J . A m . Chem. (4) N. Thorp and R. L. Scott, T H IJOURNAL, ~ 60, 670 (1966).
*
