LEOJ. PARIDON, GEORGE E. MACWOOD AND JIH-HEN*Hu
1998
Vol. 63
THE HEAT OF VAPORIZATION OF DIBORANE BY LEOJ. PARIDON, GEORGE E. MACWOOD AND JIH-HENG Hu contribution from the McPherson Chemical Laboratory of the Ohio State University, Columbus 10,Ohio Received March
4, I960
The heat of vaporization of diborane has been measured calorimetrically from the normal boiling point to within four degrees of the critical point. Over this range, the heat of vaporization as a function is given quite well by the expression Lv = 546.2 (To T)o.sg cal. mole-', where Tois the critical temperature, 289.7"K.
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Introduction As a part of this laboratory's program for the investigation of the thermodynamic properties of diborane, its heat of vaporization was determined calorimetrically from the normal boiling point to the critical point. Experimental Apparatus.-The condensed gas calorimeter used in this work is the same as that of Hu, White and J0hnston.l Procedure.-Diborane of 99.91 mo!e % purity was condensed into the calorimeter from a cahbrated volume. The uantity of material was determined to 0.4% by measuring %e pressure with a mercury manometer and a Gaertner cathetometer. All pressure readings were referred to a standard meter bar, and the observed pressures were corrected for capillary depression, the temperature of the mercury and scale, and for local value of g. The temperature of the calorimeter was determined using a standard copper-constantan thermocouple and a Wenner potentiometer. The energy input for vaporizing diborane was determined, using a double White potentiometer and a resistance thermometer. At the start of a determination, the temperatures of the calorimeter blocks, which had been adjusted previously to the temperature of the calorimeter, were recorded. The current through the resistance thermometer was read, as well as its drift rate and the galvanometer sensitivity. The potential drop across the resistance thermometer and the e.m.f. of the thermocouple were read alternately a t 30-see. intervals until steady drlft rates were obtained. The valve at the top of the calorimeter was cracked open at the same time that the heating period began. By throttling this valve it was possible to maintain the temperature of the calorimeter constant to k0.05" throughout the.heating period. The vaporized diborane was collected in .an evacuated calibrated volume. At the end of the heating period the material remaining in the lines outside the calorimeter was transferred into the calibrated volume by means of a Toepler pump. The heating period was adjusted to approximately 15 min. duration. The rate of energy input was determined by measuring the potential drop across the heater a t the midpoint of the heating period, and the current was taken as the average of that obtained a t 0.21 and 0.79 of the heating period. The energy input so obtained was corrected for heat leak as well as for the drift in temperature from the beginning to the end of the heating period.
Results The molar heat of vaporization L, is related t o the experimentally observed Q and m by2
or L" --- QM m
(I - - ;:)
where Q is the energy input necessary to vaporize m grams of diborane, M the molecular weight, p the vapor pressure, V g and V1, the molar volume of (1) J. Hu, I). White and H. L. Johnston, J . Am. Chem. Soc., 7 6 , 5642 (1953). (2) N. 8. Osborne and D. C. Ginnings, J . Research Natl. Bur. Sfandurds, 89, 453 '1947).
diborane vapor and liquid, respectively. The vapor pressure data used are reported in the preceding paper.3 The molar volumes were taken from unTABLE I HEATOF VAPORIZATION OF DIBORANE TVd
T,OK.
QM/M cal. moli-1
cal. mog-1
179.79 179.37 179.38 190.06 199.77 209.98 220.03 229.99 239.89 250.03 259.92 264.97 270.07 274.87 277.94 280.98 284.01
3412 3430 3417 3312 3198 3066 2946 2808 2664 2511 2326 2236 2121 1988 1900 1809 1646
15 15 15 25 38 56 81 112 155 207 277 320 371 418 468 531 576
LV (eq. 3), cal. mole-1
LV (ex ) cal. rnoPe'1
3397 3415 3402 3287 3160 3010 2865 2696 2509 2304 2049 1916 1750 1570 1432 1278 1070
3405 3404 3404 3286 3158 3013 2858 2692 2508 2295 2052 1908 1744 1563 1428 1271 1076
TABLE I1 HEATA N D ENTROPY OF VAPORIZATION OF DIBORANE Lv,
T, OK.
cal. mole-'
180.57 180.80 190.00 200.00 210.00 220.00 230.00 240.00 250.00 260.00 270.00 273.16 280.00 285.00 288.00 289.7
3405 3405 3287 3155 3012 2859 2691 2506 2296 2050 1747 1631 1325 999
672 0
AS", cal. mole-1 deg.-l
18.85 18.85 17.30 15.78 14.34 13.00 11.70 10.44 9.18 7.88 6.47 5.97 4.73 3.51 2.3 0
published measurements made in this Laboratory. QM were calculated using both The corrections of m
eq. 1 and 2 and were consistent. The results of the heat of vaporization measurements and calculations are given in Table I. Considering the precision of the measurements and various sources of systematic error, the maximum error (3) Leo J. Paridon and George E MacWood. THIBJOURNAL. 81. 1997 (1959).
Dee., 1959
SOLUBILITY OF HYDROGEN FLUORIDE IN NAF-ZRF~ MIXTURES
1999
measurements, reporting 3412 and 3431 cal./mole, respectively. The first value is in good agreement with the value reported here considering the uncertainty in the two values. Wirth and Palmer calculated a value of the heat of vaporization a t the L, = 546.2 ( T o - T)O.39 cal. mole-' (3) normal boiling point from their vapor pressure where T , = 289.7"K. Table 11gives the values of measurements, estimating the gas imperfection L,, calculated using eq. 3, and ASv, the entropy of using the Bertholet equation. Their value is 3413 cal./mole and is in exceptional agreement with the vaporization. There have been several previous determinations other two. (4) John T. Clarke, E. B. Rifkin and H. L. Johnston, J . Am. Chem. of the heat of vaporization of diborane at the nor- Soc., 76, 781 (1953). mal boiling p ~ i n t . ~Clarke, ,~ et al., determined it 60, 911 (5) Henry E. Wirth and Emiel D. Palmer, THIS JOURNAL, both calorimetrically and from their vapor pressure (1956).
in the molar heat of vaporization is estimated to be 10 cal./mole. To smooth the data, they were fitted to the equation
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SOLUBILITY OF HYDROGEN FLUORIDE I N MOLTEN FLUORIDES. I. I N MIXTURES OF NaF-ZrF4 BY J. H. SHAFFER, W. R. GRIMESA N D G. M. WATSON Oak Ridge National Laboratory,l Oak Ridge, Tennessee Received March 6 , 1968
Solubilities of H F have been determined a t pressures from 0.5 to 3 atmospheres over a temperature range 550 to 800" NaF. Henry's law is obeyed and the solubility decreases with inin NaF-ZrF4 mixtures containing 45 to 80.5 mole creasing temperature. The solubility of HF increases approximately tenfold as the mole % of NaF in the solvent is increased from 45 to 80.5 mole yo. Henry's law constants in moles H F per cc. of solution per atmosphere at 600, 700 and 800" are (1.23 f0.04) x 10-6 and (0.03 0.02) X !0-6 and (0.73 0.01) X 10-6 in NaF-ZrF4 (53 mole % NaF). The enthalpies of solution in kcal. per mole also increase in magnitude from -3.85 to -9.70 as the mole % of NaF is changed from 45 to 80.5. The entropies of solution vary from -5.2 to -6.5 e.u. over the same range of solvent compositions.
*
*
Introduction This investigation constitutes part of a systematic study of the solubilities of gases in molten salts. In this study an attempt is being made to elucidate the solvent characteristics which have an effect on gas solubility. The results of measurements on the solubility of H F in various mixtures of NaF-ZrF4 at pressures from 0.5 to 3 atmospheres over the temperature range 550 to 800" are presented here. Solubilities of gases in liquids have been studied by many investigators; reGews covering much of this work are Solubilities of noble gases in molten fluoride mixtures have been rep ~ r t e d . Burkhard ~ and Corbett6 recently determined the solubility of water in molten LiC1-KC1 mixtures and presented an approximate value of the solubility of HC1 a t 480" in a mixture containing 60% LiC1. No information regarding the solubilities of HF in molten fluorides in the range of pressures, temperatures and solvent compositions presented here has been found in the literature. Related investigations 011 some of the properties and phase behavior of acid fluorides of LiF,' (1) Operated for the United States Atomic Energy Commission by the Union Carbide Corporation. (2) A. E . Markham and K. A. Kobe, Chem. Reus., 28, 519 (1941). (3) J. H. Hildebrand and R. L. Scott, "The Solubility of NonElectrolytes," Chapter XV, 3rd Ed.. Reinhold Publ. Corp., New York. N. Y . , 1950. (4) M. W. Cook, U. S. Atomic Energy Commission, UCRL-2459 (1954). ( 5 ) (a) W. R. Grimes, N. V. Smith and G. M. Watson, THIS JOURNAL. 62. 862 (1958); (b) M. Blander, W. R. Grimes, N. V. Smith and G . M. Watson, ibid., 63, 1164 (1959). (6) W. J. Burkhard and J. D. Corbett, J . A m . Chem. SOL, 79, 63F1 ( 1 957).
NaF,+1° KF," RbF,12CsFi3and NHI FI4have been reported. Information on liquid-liquid solubilities and liquid-vapor equilibrium of HF-UFe mixtures is available.15p16 Mixtures of NaF and ZrF4 were chosen for this investigation because of their value in nuclear fuel element r e p r o c e s ~ i n g ~and ~ - ~their ~ use (with added UF,) as fuel for an experimental nuclear reactor. 22 (7) H. V. Wartenberg and 0. Bosse, Z . Eleklvochem., 28, 386 (1922). (8) J. F. Froning, M. K. Richards, T. W. Stricklin and S. G. Turnbull, Ind. Eng. Chem., 39, 275 (1947). (9) D. G. Hill, unpublished work, Oak Ridge National Laboratory, Oak Ridge, Tennessee. (10) W. Davis, Jr., KLI-2552, "Vapor Pressure-Temperature Relations in the System NaF-HF," Union Carbide Nuclear Company, Oak Ridge Gaseous Diffusion Plant. ( 1 1 ) G. H. Cady, J . Am. Chem. SOC.,66, 1431 (1934). (12) E. B. R. Prideaux and I
