Sept., 1959 Acknowledgment.-This work was reported under Contract AF 33(616)-6057 with Wright Air Development Center. Their permission to publish is gratefully acknowledged.
THE HEAT OF VAPORIZATION AND VAPOR PRESSURE O F CARBON TETRACHLORIDE; THE ENTROPY FROM CALORIMETRIC DATA BYD. L. HILDENBRAND AND R. A. MCDONALD Thermal Laboratory, The Dow Chemical Company, Midland, Michigan Received February 8, 1969
Calculated values of the thermodynamic properties of the highly halogenated methanes are somewhat uncertain because the large vibrational contribution makes it important that the frequencies be known rather accurately. I n addition, it is necessary to include significant corrections for vibrational anharmonicity, the data for which are seldom available. This is particularly true of carbon tetrachloride. For these reasons, i t is important that the calculated thermodynamic properties be checked against accurate calorimetric vapor heat capacity and entropy data. It is the purpose of this note to summarize the best available calorimetric entropy data for carbon tetrachloride by combining newly measured heat of vaporization and vapor pressure data with available low temperature heat capacity data. Hicks, Hooley and Stephenson1 have reported low temperature thermal data for CC1, covering the range 17 to 298°K. Their results appear to be the most reliable and complete of a number of low temperature investigations. The recently reported liquid heat capacity data of Harrison and Moelwyn-Hughes2 agree with those of Hicks, Hooley and Stephenson1 within the 0.5% accuracy claimed by the former authors. Based on JCpd In T , Hicks, Hooley and Stephenson1 have reported the entropy of liquid C C 4 a t 25" as 51.25 f 0.15 cal. moleA1deg.-l. However, there appears to be a serious error in the extrapolation of the heat capacity below 18°K. Hicks, Hooley and Stephenson' made the extrapolation using a Debye function with 0 = 148" and 15.4 degrees of freedom, obtaining an entropy of 1.39 cal. mole-l deg.-l a t 18°K. Experience in this Laboratory has shown that the empirical use of Debye functions with large numbers of degrees of freedom leads to significantly lower heat capacities and entropies than would be obtained with a Debye function of three degrees freedom plus the appropriate number of Einstein functions. This latter combination of functions is believed to give an extrapolation that is more acceptable from a fundamental view, since it allows three Debye degrees of freedom to account for the lattice contribution and makes up for internal vibrational degrees of freedom with Einstein functions. Aside from this, however, the extrapolation of H., H. and S. fits only the first three C, (1) J . F. G. Hicks, J . G. Hooley and C. C . Stephenson, J . A m . Chem. Soc., 66, 1064 (1944). (2) D. Harrison and E. A. Moelwyn-Hughes, Pvoc. Roy. Soc. (London), mas, 230 (1957).
KOTES
1521
values and above 20°K. it diverges rapidly from the experimental points. On the other hand, a Debye function (three degrees freedom) with 0 = 70" together with four einstein functions each with e = 98" gives an excellent fit over the range 17 to 32°K. and leads to an entropy of 1.805 cal. molev1 deg.-' a t 18°K. This latter extrapolation has been adopted for the entropy calculation and, when combined with the graphical integration of H., H. and S., gives an entropy of 51.67 f 0.15 cal. mole-l deg.-l for liquid CCl4 a t 25'. The entropy of ideal gaseous CCL can be evaluated from the above data and the heat of vaporization and vapor pressure data reported here. The calculation is summarized in Table I. The Berthelot equation and the critical constants given by Kobe and Lynn3 (To = 556.4'K., Pc = 45.0 atm.) were used in evaluating the gas imperfection correction. TABLE I THEENTROPY OF CARBONTETRACHLORIDE, CAL. MOLE-' DEC.-l 51.67 k 0 . 1 5 Satd. liquid, 298.15"II. Vaporization, 7746/298.15 25.980 Gas imperfection 0.036 Compression, R In (114.2/760) -3.766 Ideal gas, 298.15'K.
73.92k0.20
For comparison with our calorimetric value, the most reliable calculated entropy appears to be that given by Albright, Galegar and Innes14 which is based on the infrared data of Plyler and B e n e d i ~ t , ~ the Raman measurements of Claassen6 and includes a semi-empirical anharmonicity correction. The anharmonicity correction was obtained by correlation with the calculated and observed gas heat capacities of CC12F2, and is considered to be only an approximation. Interpolation in the tables of Albright, Galegar and Innes4 gives an entropy of 74.12 cal. mole-l deg.-l a t 298.15"K., which agrees with the calorimetric value within experimental error. The calculated value is probably uncertain by about 0.1 cal. mole-l deg.-l. Gelles and Pitzer' give a calculated entropy of 73.94 cal. mole-1 deg.-l a t 298.15"K., in good agreement with the calorimetric value, but their value is based on only the infrared measurements of Plyler and Benedict6 and does not include an anharmonicity correction. Accurate vapor heat capacity measuremeiits would be very helpful in further reducing uncertainties in the calcuhted thermodynamic properties. Experimental Material.-Commercial carbon tetrachloride was purified by fractional distillation. A center cut was found to have it purity of 99.96 mole % by freezing curve analysis. Heat of Vaporization.-The heat of vaporization of CC1, at 25' waa measured using a newly constructed calorimeter that will be described in detail in a later publication. Briefly, (3) K. A. Kobe and R . E. Lynn, Chem. Reus., 6 9 , 117 (1953). (4) L. F. Albrigbt, W. C. Galegar and K. K. Innes, J . Ant. Ckem. Soc., 76,6017 (1954). ( 5 ) E. K. Plyler and W. S. Benedict, J . Research Natl. Bur. Standards. 4T,202 (1951). (0) H. H. Claassen, J . Chem. Pliys., aa, 50 (1954). (7) E. Gelles and K. 8. Pitzer, J . A m . Chem. Soc., T6, 5259 (1953).
1522
NOTES
the system consisted of a calorimeter vessel with glass exit tube mounted within a conventional adiabatic shield assembly and insulated by vacuum. The adiabatic control system is the same as that used on the low temperature calorimeter in this Laboratory.* The cylindrical sample container was constructed of copper and was fitted with a central re-entrant well containing a platinum resistance thermometer (Ro= 91 ohms), an internal heater wrapped around the base of the thermometer well and a system of 40 radial heat transfer vanes. The heater was wrapped over only the lower one-fourth of the thermometer well so as to avoid superheating the vapor leaving the calorimeter. All parts of the sample container were nickel-plated before assembly. The calorimeter had a volume of 60 0111.3 The temperature of the vapor leaving the calorimeter was measured with a chromel-constantan thermocouple attached to the short platinum exit tube soldered to the calorimeter top and could be maintained constant within 0.05” by manipulation of a throttle valve in the exit line. The exit line was provided with electrical heating where necessary in order to prevent condensation. Material vaporized was condensed in liquid nitrogen cooled glass bulbs and weighed. Corrections were applied for the heat effect corresponding to small changes in the temperature of the calorimeter and for the heat required to vaporize material replacing the liquid withdrawn from the calorimeter. As a check on the operation of the calorimeter, the heat of vaporization of a sample of high purity benzene was measured at several temperatures and compared with the accurate results of Osborne and Ginning@ and Waddington and Douslin.10 The results of seven vaporization experiments over the range 20 to 60” all agreed with the accepted values to within O.l%, the average difference being 0.04%. Varyin the rate of vaporization produced no observable effect. %e measurements on CC14 are summarized in Table 11. Uncertainties in the experimental value should not exceed 0.1%.
Vol. 63
from the measured boiling temperatuyes and the known vapor pressure of water.9 The experimental points are given in Table 111. The constants in the Antoine equation 227.16) log P(mm.) = 6.89406 - 1219.58/(t were obtained by a least squares treatment. Values calculated from the above equation are compared with the experimental points in Table 111.1’
+
(11) N. S. Osborne. H. F. Stiinson and D . C. Ginnings, J . Reaearch Natl. BUT.Slandavds, 28, 261 (1939).
-
THE HEAT OF FORMATION OF METHYL NITRATE BY JAMESD. RAY‘A N D R. A. OGG,JR. Contribution /?om the Department o f Chemistry, Stanford University, California Received February 9, 1969
Recent interest in the kinetics of reactions involving methyl nitrate led the authors to a calorimetric investigation of the gas phase heat of the reaction between nitrogen pentoxide and methyl nitrite to form an equilibrium mixture of nitrogen dioxide and dinitrogeii tetroxide plus methyl nitrate. Preliminary studies by infrared analysis showed that the products of the reaction of equal pressures of reactants represented quantitative conversion with no evidence of side reactions. The reaction was found to be extremely rapid.
Materials.-Methyl nitrite was prepared by the reaction TABLE I1 given in “Organic Syntheses.”* The reaction was carried THEHEATOF VAPORIZATION OF CARBON TETRACHLORIDE out in a vacuum system. The product was purified by condensation on anhydrous potassium carbonate and by disAT 25’, CAL.MOLE-‘ tilling repeatedly from Dry Ice to liquid nitrogen tempera1 , “C. G . , vaporized Rate g. m h - 1 AH AH?sa tures. An infrared spectrum of the product showed no 24.17 8.340G 0.278 7757 7747 bands of possible impurities such as nitric acid, nitrogen dioxide or methyl nitrate. Nitrogen pentoxide of at least 24.68 8.3857 .279 7746 7742 99.9% purity was prepared as previously described.* 24.58 5.2625 .175 7754 7749 Apparatus and Experimental Procedure .-The calorimetcr Accepted value = 7746 i 5 employed has been described previous1y.a It contained a 275-m1. gas reaction bottle. The heat capacity of the enFor comparison purposes, the hcat of vaporization at 25’ tire calorimeter including the chlorobenzene li nid was was calculated from the exact Clapeyron e uation, using the 168.4 f 1.1 cal./deg. In the prcsent study, afterlollowing Antoine vapor pressure equation given baow and the Ber- the iaitial temperature drift, 75 mm. pressure of nitrogen thelot equation to obtain the vapor volume. The value thus pentoxide was admitted to the calorimeter bottle. This was calculated is 7730 cal. mole-’, in reasonable agreement as high a pressure as could be used with no question as to with the experimental value of 7746 cal. mole-’. This absence of a liquid phase forming from the products. Onewould indicate that the use of the Berthelot equation to half minute later, methyl nitrite was admitted quickly to a obtain the gas irnpcrfection correction is reasonable. pressure of 248 mm., which corresponded to a 69.5 mm. Vapor Pressure.-Measurements of the vapor pressure final excess. I t was shown by Yoffe and Gray4 that t’ie nitrowere made using a twin ebulliometer system, with water gen dioxide product does not react with this excess of as thc reference substance. Pressures were determined methyl nitritc. The temperature drift was then again followed. The final temperature was 25’. The system W ~ Z R TABLE 111 thermostated at the mean reaction temperature. Thc ohserved tempertture rise (average of 4 determinations) was T H E VAPOR PItESBURE O F CARBON TETRACHLORIDE 0.070 =t0.002 . The equilibrium pressures of nitrogen diP(obsd.) - P(oa1cd.) oxide and dinitrogen tetroxide present in the calorimeter t . oc. P(obsd.) mm. mm. bottle after the reaction were 57.0 and 46.5 mm., respec19.88 00.60 -0.03 tively. The stoichiometric equation for the reaction wlth 36.93 189.10 .20 the standard heat of formation of each compound from the elements written below each respective tcrm is 56.16 387.97 - .50 75.28 7G. 16 7G.87 77.71 (76.73)
727.30 747.70 763.03 782.83 (760.00)
+ + + -
.20 .75 .24 .I1
....
+
1.000 Nz05 1.000 CHyONO = +3.17 -1G.27 0.620 NzOi 0.760 NO2 $0.620 X 2.239 +760 X 7.9G4
+
+ 1.000 C H I O N O ~ AHr(CH30NOd
( 1 ) Department of Chemistry, University of Minnesota, Minne-
(8) D. L. Hildebrand, W. R. Kramer, R. A. McDonald and D . R. 6tull. J. Am. Chem. Sor., 80, 4129 (1958). (9) N.8. Osborne and D. C. Ginnings, J . Research Natl. Bur. Standards, 39,453 (1947). (10) G . Waddington and D. R. Douslin, J . A m . Chem. Soc., 69, 2275 (1947).
apolis 14, Minn. (2) W. H. Hartung and F. Crossley, “Organic Syntheses,” Coll. Vol. 11, A. H. Blatt, editor, John Wiley and Sons, Ino., New York, N. Y., 1943, p. 363. (3) R. A. Ogg, Jr., and J. D . Ray, THIS JOURNAL, 61, 1087 (1957). (4) A. D . Yoffe and P. Gray, J . Chem. SOC.,1412 (1951).
L
