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PHYSICAL CHEMISTRY Registered in U . S. Patent Ofice @ Copyright, 1969, by the American Chemical society
Thermodynamics of Iodine Trichloride, Entropy and Heat Capacity from 15 to 325'K. Composition of the Equilibrium IC1 and C1, Gas Phase1 by R. H. Lamoreaux and W. F. Giauque Low Temperature Laboratory, Departments of Chemistry and Chemical Engineering, University of California, Berkeley, California 94720 (Received October 1 4 . 1 9 6 8 )
The heat capacity of IC13 has been measured from 13 to 325°K. Its entropy was found to be 40.40 gibbs/mol and its free energy and heat of formation from the elements, AP" = -5413 cal/mol, AHo = -21,340 cal/mol, each at 298.15"K. AH"o = -21,723 cal/mol. Values of Cpo, X", (Po IIoo)/T, and ( H o - Hoo)/T have been tabulated. The temperature of the ICla-IC1 eutectic has been found to be 296.21"K and the partial pressures over the eutectic are P(IC1) = 23.9 Torr and P(C12) = 19.7 Torr. For the reaction ICla(s) = ICl(g) Cln(g) , AHoo = 26,297 cal/mol. The vapor over ICla(s) and the equilibrium liquid on the IC1 side is known to consist of IC1 and C4. Their partial pressures have been evaluated over the range 296.56312.78"K. Below 301.5'K, IC1 predominates in the vapor and above this temperature an excess of Cla increases rapidly. An error in the calculations of Calder and Giauque, J. Phys. Chem., 69, 2443 (1965), has been noted for the reaction ICl(s) (1) = ICl(g). AHOo should have been 13,245 cal/mol instead of 13,273 cal/mol used in their calculations. Revised values of the vapor pressure of IC1 and data relating to AP and A E for several reactions are given in an addendum to the present paper.
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This paper reports a continuation of the work in this laboratory on the thermodynamics of the iodinechlorine system. A previous paper2 has reported heat capacity data to low temperatures on iodine monochloride together with measurements of the composition and pressure of the equilibrium vapor over the liquid. The present paper reports heat capacity data on IC18 over the range 15-320°Kand measurements of the Cl/I atomic composition ratio in the gas over solid IC18 and the liquid which is in equilibrium with it on the IC1 side of its melting curve. The above ratio a t the ICla:ICl eutectic temperature, which was evaluated from the measurements, enabled the calculation of the partial pressures of IC1 and Clz gases over this solid-liquid equilibrium system. These data, combined with the previous work on IClJ2determine the free energy of formation of IC&,which, with the entropy values based on the third law of thermodynamics, enables evaluation of its heat of formation from IC1 and Cla or from 312 and jC&. Previous measurements on this system, including the
above Cl/I composition ratio in the gas, have been made by Stortenbeker3p4and by Nies and Yoste6 The present measurements of the Cl/I composition ratio should be considerably more accurate due to the use of constant stirring in the thermostated solid-liquid system, and the taking of large samples for analysis through a fine capillary over a very long time period, so as not to disturb the equilibrium.
Experimental Section Calorimetric Apparatus. The calorimetric apparatus of Giauque and Egan6was used except that a Hastelloy Type C alloy' calorimeter was substituted for the (1) This work was supported in part by National Science Foundation Grant GP-8137. (2) G. V. Calder and W. F. Giauque. J. Phys. Chem., 69, 2443 (1965). (3) W. Stortenbeker, Rec. Trav. Chdm., 7 , 152 (1888). (4) W. Stortenbeker, 2.Phys6k. Chem., 3 , 11 (1889); 10, 183 (1893). ( 5 ) N . P. Nies and D. M . Yost, J. Amer. Chem. Soc., 5 7 , 306 (1935). (6) W. F. Giauque and C. J. Egan, J. Chem. Phys., 5 , 45 (1937). (7) Union Carbide Co., Kokomo. Ind.
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R. H. LAMOREAUX AND W. F. G I A C Q U ~
original gold calorimeter to avoid corrosion. The ability of Hastelloy to successfully resist corrosion in this interhalogen system had been investigated in connection with the work on IC1.2 A sample of IC13 was kept in the Hastelloy shell of the final calorimeter for 20 days and then removed. The vessel was rinsed with aqueous KI, then with distilled water, and baked at 110" while being flushed with dry iS2 gas. The thermal conductivity of Hastelloy is so low that it is far from an ideal material to assist the attainment of equilibrium in a calorimeter. In an attempt to improve this situation, we tried the suggestion of Calder and GiauqueJ2who suggested the addition of a copper layer made from a tube soldered to the exterior of the Hastelloy. The idea might have succeeded if the copper had been welded to the Hastelloy; however, the solder could not resist the differential contraction on cooling, leading t o breakage of the gold thermometerheater wound on the exterior of the calorimeter, which was discarded. To avoid this difficulty, a second Hastelloy calorimeter was enclosed in a slightly loose copper jacket, onto which the gold thermometer-heater was wound. The space between the Hastelloy and copper was filled with helium gas. The arrangement was a substantial improvement over the original Hastelloy calorimeter. The completed calorimeter is shown in Figure 1. Hastelloy C tubes, 0.254 cm 0.d.
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Giauquee2 In addition, sections A-A and B-B show flattened Hastelloy tubes, 5, which were welded into the cylindrical calorimeter as indicated. These 0.475-cm diameter tube sections were 6.5 cm long with a 0.013-cm wall. They were rolled nearly flat so that copper strips 5.5 cm long and 0.013 cm thick could be inserted before the final rolling. The Hastelloy ends were sealed with an atomic hydrogen torch before they were formed into semicircles and electrically spot-welded in place. Thirteen pairs of vanes, arranged in a staggered configuration, were placed in the calorimeter to improve lateral thermal conductivity. The volume of the calorimeter was 144 cma. It seems probable that a very small amount of hydrogen gas was enclosed in these tubes during the welding operation. This would assist heat transfer between the copper strips and their enclosing Hastelloy jackets. The outer cylindrical copper jacket, 6, was added to provide longitudinal thermal conductivity, after the calorimeter had been filled with IC&. It was desirable not to have longitudinal conductivity during the filling operation to be described later. Epoxy joints, 2,2, were used for the final closure of the copper jacket to avoid the heat required for soldered joints, A detachable thermocouple well is shown at 7. This device has been deEicribed by Busey and Giauque.8 A gold thermometer-heater constructed as described by Murch and GiauqueQwas wound on the exterior of the copper jacket. Standard copper-constantan thermocouple No. W-23 was used as a temperature reference. It was compared with normal hydrogen triple point, 13.95'K, and boiling point, 20.37'K; and nitrogen triple point, 63.15"K, and boiling point, 77.33"K. Appropriate corrections were made for the small change since the original calibration in 1933. 0OC was taken as 273.15"K. One defined calorie was taken as 4.18400 absolute joules. Preparation of IC&. The general procedure was similar to that described by Calder and Giauque2 for the preparation of IC1. The chlorine was purified in a low-temperature vacuum-jacketed silvered fractionation column as described previously.2 The Mallinckrodt iodine used was stated to contain only 0.005% C1 Br and 0.01% of nonvolatile material. All analyses were made by weighing the combined halogens followed by conversion to 13-ion with aqueous K I solution and weight titration with aqueous NazSzOa. Filling the Calorimeter. According to St~rtenbeker,~ ICls "melts" only at some temperature near 101", and at a pressure of some 16 atm.3 The problem of introducing such a material to a calorimeter is somewhat complex. An attempt was made with the first calorimeter by
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Figure 1. Hastelloy calorimeter, with internal Hastelloy jacketed copper heat conductors and external copper jacket, for use with halogens.
and 0.17 cm i.d., are shown as 1 and 3. They were used in the procedure for filling the calorimeter with IC&, as described later. The Hastelloy C calorimeter is indicated by 4. The dimensions of the Hastelloy shell were identical with those given by Calder and The Journal
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( 8 ) R. H. Busey and W. F. Giauque, J. Amer. Chem. Soc., 74, 4443 (1952). (9) L. E. Murch and W. F. Giauque, J. P h y s . Chem., 6 6 , 2052
(1962).
THERMODYNAMICS OF IODINE TRICHLORIDE adding a known amount of the Iz-ICl eutectic as liquid, because of its low vapor pressure. This was followed by the condensation of an amount of liquid chlorine which was slightly less than the amount required to form stoichiometric IC&. After sufficient cooling and the addition of some helium for thermal conduction, the filling tube of the Hastelloy calorimeter was closed by pinching it, followed by heliarc welding. The calorimeter was then heated to 110"to facilitate reaction. Some preliminary heat capacity measurements showed heat effects near the melting point of chlorine and were taken as a definite indication that equilibrium combination had not occurred. It is probable that the helium interferred .with the reaction since the necessary diffusion would be slow. Also, since the substance is considerably decomposed into Clz and IC1 gases near lOO", with a substantial excess of Cl2, stoichiometric recombination with the resultant IC1 rich liquid may have been difficult in the range below 100". We visualize an IC1-rich liquid at the bottom, the IC18 solid intermediate, with the Clt rich gas above. Although the calorimeter was repeatedly turned end over end during the cooling, complete recombination was not attained. An alternate procedure was then substituted.
Figure 2. Schematic diagram of apparatus used for filling calorimeter with solid iodine trichloride by using chlorine as a carrier gas.
Figure 2 is a schematic drawing of the arrangement used for condensing the solid IC13 within the second Hastelloy calorimeter, 10, of Figure 2. The Hastelloy tubes, 8 and 9 of Figure 2, were sealed to common glass which was attached to the Pyrex apparatus by means of graded seals. A 1:3 mixture of KeLF200 wax with Kel-F210 greaselo was used2 to lubricate stopcocks.
757 The method of transfer recognizes the fact that solid IC18 decomposes to give only IC1 and Cl2 in the vapor. This was shown by the early but accurate gas density experiment of Melikoff The present experiments show that above 301.5"K the vapor contains a C12/IC1molecular ratio in excess of unity. At first it was thought that a simple circulating system, with the lower part of the calorimeter cooled to the ice point and the IC& supply flask, 4 of Figure 2, and other upper portions of tJhesystem heated to about 333°K would transport the material to the calorimeter by simple convection, using the Clz as a temporary carrier gas. The Hastelloy inlet and outlet tubes of the calorimeter, 8 and 9 in Figure 2, were heated electrically. This method was undoubtedly correct in principle but too slow. The method was then modified with the assistance of stopcocks 2, 5 , 6, and 7 and a side tube, 3, which could be cooled to 196"K, the temperature of solid COZ. At the ice point the vapor pressure over IC13 is quite low and the partial pressure of IC1 exceeds that of Cl2 unless Clz is present in excess. The vapor pressure of liquid chlorine at 196°K is too high to remove chlorine from IC13 at 273°K by decomposition. At 333°K the vapor pressure is of the order of 1 atm and the CL partial pressure is greatly in excess of the IC1 pressure. The Hastelloy calorimeter was filled before the copper jacket was added; thus it had quite poor longitudinal conductivity and the lower portions could easily be maintained near the ice point while the top and tubes 8 and 9 remained warm. The sequence of the filling operation was as follows. (A) Stopcocks 5,6, and 7 were open and stopcock 2 was closed. The side tube, 3, was at 196"K, the supply bulb, 4,was at 333"K, and the bottom of the calorimeter, 10, was a t 273°K. The condensation of liquid Clz in the side tube greatly accelerated the transfer. However, the removal of Clzeventually depleted the chlorine fraction of the original IC13 in supply bulb 4 and the process slowed. (B) Stopcocks 5 and 6 were closed and 2 was opened. Side tube 3 was warmed and supply bulb 4 cooled somewhat until the chlorine was transferred to it. Sequence A was then repeated. After three cycles, the IC13 was nearly all condensed in the calorimeter as a solid. Stopcocks 6 and 7 were then closed and 2 was opened so that all residual amounts of halogen in side tube 3 and supply bulb 4 could be removed through stopcock 1 for analysis. The h a l result showed a small deficiency of chlorine, which was considered preferable to an excess, since data were available2 to correct the heat capacity for a small amount of solid IC1 obtained as eutectic mixture in the cooled sample. The calorimeter contained 0.99922 mol (10) Minnesota Mining and Manufacturing C o . , St. Paul, Minn. (11) P. Melikoff, Rer. Deut. Chem. Ges. (Berlin), 8 , 490 (1875). V0'olum.eYS,Number 4 April 1060
R. H. LAMOREAUX AND W, F. GIAUQUE
758 of IC& and the chlorine deficiency produced 0.00584 mol of IC1 in the solid sample investigated. The calorimeter was then cooled in ice and an atmosphere of helium was added. Tubes 1 and 3 of Figure 1 were then pinched off and sealed by heliarc welding. The copper jacket, 6, of Figure 1, was then added. T h e Eutectic Temperature of Iodine Monochloride and Iodine Trichloride. Since the eutectic equilibrium gives an important relation between the IC1 and ICls systems, its temperature was determined in separate experiments. IC& (0.65 mol) and IC1 (1.60 mol) were added to a closed Pyrex glass bulb with a well for a copperconstantan thermocouple. Thermal contact between the well and thermocouple junction was obtained by means of kerosene. After cooling the assembly to about O", it was allowed to warm slowly and the temperature stayed constant for a long period (days) at 23.06" = 296.21 f 0.01'K. The experiment was repeated with an identical result. Heat Capacity Observations. The heat capacity data are recorded in Table I. As mentioned above, the sample was slightly deficient with respect to chlorine
Table I: Heat Capacity of Iodine Trichloride Sarnplea,*,C T,,, 'I
IZ(S)(C & (2-23)
A . F ' z ~ ~ . ~=s ~3314 K f 25 cal/mol
A Thermodynamic Study of the Urania-Uranium System1 by R. J. Ackermann, E. G. Rauh, Chemistry Division, Argonne National Laboratory, Argonne, Illinois
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and M. S. Chandrasekharaiah2 Chemical Engineering Division, Argonne National Laboratory, Argonne, Illinois
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(Received May 9 , 1 9 6 8 )
The phase diagram (temperature-pressure-composition) for the uranium-oxygen system from 1600 to 2500°K in the composition range of the unvariant system UO,(s) U(1) was investigated in three steps. The composition of the urania phase in equilibrium with liquid uranium from 1900 to 2500'K was determined by an isopiestic method. The mole fraction of uranium dissolved in the UOZ phase is given by log N U = (1.63 f 0.11) - (6270 f 240)/T. The total effusion rates were measured by means of the effusion method from 1580 to 2400'K. iin "effective" vapor pressure p , was calculated from the measured effusion rates, assuming the vapor to be UO(g). The least-squares treatment of the data resulted in the expression, log p,(atm) = (7.25 f 0.15) - (27,020 f 250)/T. The mass spectrometric investigation of the same system has shown the vapor to be predominantly UO(g) with smaller amounts of U(g) and UO,(g). The temperature dependence and relative intensities of the species were measured mass spectrometrically and combined with the total effusion rates to yield the partial pressures for each species: log pu(atm) = (5.21 & 0.14) - (25,640 f 300)/T; log puo(atm) = (7.11 f 0.14) - (26,880 & 300)/T; log puo,(atm) = (7.74 f 0.14) - (30,180 =t300)/T. The standard free energy of formation of UO(g) is given by the equation AG?(UO) = -7800 - 13.82' (kcal/mol)
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I. Introduction
It is now generally recognized that many so-called chemical compounds can exist as phases of variable composition.g Uranium dioxide (urania) is certainly no exception and, in fact, experimental evidence suggests that its composition can vary continuously within wide limits. The crystallographic and thermodynamic properties of the superstoichiometric phase which extends up to a composition given approximately 2.2, have been by the oxygen-to-uranium ratio, O/U most extensively investigated in recent years.4 Early evidence which suggests the existence of a substoichiometric urania phase which may extend down to O/U = 1.75 was compiled by R r e ~ e r , but ~ more recently others including Anderson, Sawyer, Worner, Willis, and Bannister,e Rothwell,' Aitken, Brassfield, and Fryxel1,s Martin and E d w a r d ~ ,and ~ Ackermann, The Journal of Physical Chemistry
Rauh, and Chandrasekharaiahl" have demonstrated the stability of substoichiometric compositions at high (1) Based on work performed under tho auspices of the U. S. Atomic Energy Commission. (2) Chemistry Division, Bhabha Atomic Research Centre, Tronibay, Bombay 74, India. (3) R. F. Gould, Ed., "Non-stoichiometric Compounds," Advances in Ohemistry Series, American Chemical Society, Washington, D. C . , 1963. (4) Report of the Panel on "Thermodynamic and Transport Properties of Uranium Dioxide and Related Phases," Technical Reports Series No. 39, International Atomic Energy Agency, Vienna, 1965. (5) L. Brewer, Chem. Rev., 52, 1 (1953). (6) J. S. Anderson, J. 0. Sawyer, H. W. Worner, G. M. Willis, and M. J. Bannister, Nature, 185, 915 (1960). (7) E . Rothwell, J . Nucl. Mat., 6, 229 (1962). (8) E . A. Aitken, H. 0. Brassfleld, and R. E . Fryxell, "Thermodynamics," Vol. 11, International Atomic Energy Agency, Vienna, 1965,p 435. (9) A. E. Martin and R . K. Edwards, J . Phys. Chem., 69, 1788 (1965). (10) R. J. Ackermann, E. G. Rauh, and M. 9. Chandrasekharaiah, Argonne National Laboratory Report, ANL-7048, July 1965.
