PHASE EQUILIBRIA OF THE SYSTEM CIFs-HF
August, 1957
1101
PHASE EQUILIBRIA OF THE SYSTEM CHLORINE TRIFLUORIDEHYDROGEN FLUORIDE' BY R. M. MCGILL,W. S. WENDOLKOWSKI, AND E. J. BARBER Union Carbide Nuclear Company, Oak Ridge Gaseous Diffusion Plant, Oak Ridge, Tennessee Received April I f 1967
The solid-liquid and liquid-vapor hase equilibria of the s stem chlorine trifluoride-hydrogen fluoride have been inyestigated over the composition range o f 0 to 100 formula % &Fa. The solid-liquid phase diagram is simple with a single eutectic occurring a t -88.8' and 7.5 formula % CIFa. The phenomenon of enantiotropy in chlorine trifluoride was observed at -83.1'. Liquid-vapor equilibria were determlned a t total pressures of 1.5, 2.0, 2.5, 3.0 and 4.0 atmospheres absolute. by means of a nickel Othmer-type equilibrium still. The system is azeotropic at all these pressures, the minimum boiling azeotropic mixture containing about 67 formula % CIFa. There is no indication of the formation of a stable complex.
Introduction Investigation of the phase equilibria of the system chlorine trifluoride-hydrogen fluoride was undertaken as a part of a series of studies of the basic chemical and physical properties of the interhalogens. The behavior of the system is of theoretical interest since the mixture is composed of the strongly polar, highly associated hydrogen fluoride which is capable of hydrogen bonding and the comparatively weakly polar and weakly associated chlorine trifluoride. Such a system is predicted to show large deviations from ideality and to have a very considerable excess free energy of mixing. A qualitative indication of the strength of bonding between chlorine trifluoride and hydrogen fluoride may be obtained by noting the presence or absence of compound formation in the solid-liquid equilibrium diagram and the degree of deviation from Raoult's law.
Experimental Phase Components.-The chlorine trifluoride was purified by passage over sodium fluoride and distillation in a twenty plate, nickel column packed with 1/8-inch nickel helices. The distillate which was employed in these studies boiIed at 11.75' at 760 mm. pressure, melted a t -76.34", and was round to be 99.96 mole yo chlorine trifluoride by thermal :innlysis (assuming monomolecular CIFa). Commercial hydrogen fluoride was purified by a teclinique developed by Rosen and Wendolkowski and described by Jarry and Davis.* By using chlorine trifluoride t o pretreat all transfer systems and receivers, hydrogen fluoride of 99.9 mole % purity (determined by thermal analysis assuming monomolecular HF) could be prepared and preserved. Solid-Liquid Equilibria.-An all-metal freezing point cell (Fig. 1 ) with a sealed-in magnetic stirrer to ensure homogeneity was used to determine the solid-liquid equilibrium. Welded nickel was employed throughout the cell except for the Monel plunger housing. The heat exchanger was provided to prevent excess conduction of heat from the solenoids to the sample chamber via the stirrer shaft housing whose free volume above the sample pot was kept at a minimum. The sample pot was suspended in a 1.5-in. diameter well in a cylindrical copper shield mounted in a double Dewar flask arrangement as shown by Rutledge, Jarry and Davis.* Temperature measurements of both the shield and sample were made using copper-constantan thermocouples calibrated with a platinum resistance thermometer over a temperature range of -78 t o +70". The over-all accuracy of the calibration was h 0 . 3 " . Arrangement of the thermocouples was such that the electromotive force could be measured either by a single point recording (1) This work was done at the Oak Ridge Gaseous Diffusion Plant operated for the government by Union Carbide and Carbon Corporation. (2) R . L. Jarry and W. Davis, Jr., T ~ r JOURNAL, s 57,600 (1953). (3) a. P. Rutledge, R. L. Jarry and W. Davis, Jr., ibid., 57, 541 (1953).
llicromax equipped with a switching mechanism or by a White double potentiometer. llicromax tracings were used to scan the various temperature regionsfor phase transitions, then the accurate transition temperature measurements were made on the White double potentiometer. Individual mixtures were made up by weight in a thin wall nickel tube easily handled on an analytical balance. T1.e hydrogen fluoride was admitted first and the chlorine trifluoride \vag condensed upon it, weights being taken at appropriate times during the procedure. The mixture of knoivii composition was transferred vapor-wise into the freezing point cell. After assembly into the temperature control equipment, the stirrer was activated and its mot/ori sustained as long as possible. For practically all warming and cooling curves, the heat transfer rate used was such a6 to give a temperature change of 0.17" per minute. For cooling curves, liquid nitrogen was placed in the outermost Dewar flask and the cooling rate controlled by changing the gas pressure in the inner Dewar flask. \\'hen the cooling cycle was complete, the liquid nitrogen wae removed, and the shield temperature was allowed to rise uniformly. Liquid-Vapor Equilibria. Equilibrium Still.-The liquidvapor equilibria data were obtained using an eyuilib1:iunl still which was similar in size and design to that described by Barber and Cady,' but was of all-nickel welded construction wherever ieasible. The still w-as mounted in a thermostated box. The pressure was determined by means of a Booth-Cromcr transmitter6 balanced by a relay mounted i n a nitrogen pressured container. The inert gas pressure was measured on either A mercury differential manometer or 3 \Vallate and Tiernan pressure gagr having :i pressure range of 0 t o 508 (mi. The hoiling ternperaturo wad measured by means ot' a calibrated tivr junrtion copper-colistantan thcrniopile and a type I\ were provided to permit simultaneous renioval of 1-rnl . portion? oi the pot charge and its cquilibriwn condeiisatc. The vacuum-jaclteted Inoilei. tvhic.11 had an opei ating liquid volume of 25 ml. was filled through the chargiiig valve at the top. The operation of the still could be follon.ed by observing the liquid condensute level arid the flow rate through a chlorofluoroplastic wiiidow in the condensate block. Attached to the still \vas a large ballast volume which could be bled through a soda-lime trap to protect the mechanical pump and the oil diffusion pump from damage by the phase components. A liquid inert to the action of chlorine trifluoride and hydrogen fluoride was circulated through the cooling condenser. To facilitate alteration of the boiler composition, a small metering manifold consisting of a 30-m1. nickel storage cylinder of hvdrogen fluoride and a graduated chlorofluoroplastic messur'ing tube \vas attached to the charging valve. Procedure.-The still n a s charged with purr chlorine trifluoride by liquid transfer. The entire composition range was then studied at each pressure b,v metering hydrogen fluoride into the still charge after each run to compensate for the volume of charge lost by sampling. Thus, if two milliliters was withdraivn in a run, two milliliters of pure hydrogen fluoride was added for the next run. The low formula weight of hydrogen fluoride and emall size of the (4) E. J. Barber and G. 1%. Cady, J . Am. Cliern. Soc., 75, 4347 (1951). (5) S. Cromer,
U.S.A.E.C. Declassified Report MDDC-803.
1102
R. M.MCGILL,W. S. WENDOLKOWSKI AND E. J. BARBER ,ANCHOR
M O N E L PLUNGER
the liquid had disappeared, a stream of helium flowing a t a rate of about 100 cc./min. was passed through the sample tube and trap. Flushing was continued until gas exhausted from the trap no longer discolored a moistened potassium iodide-starch test paper. The weight gain of the trap was taken as the weight of hydrogen fluoride in the sample; the weight loss of the sample tube was the total sample weight. From these data the composition of the samples was calculated. When the total amount of hydrogen fluoride absorbed by the sodium fluoride was approximately half of that required to form the compound NaFVHF, the charge of sodium fluoride was regenerated by heating to 400' in vucuo. Tests of this procedure made on a series of synthetic mixtures of known coniposition indicated an accuracy of 0.5 formula % over the range 0 to 90 formula % ' chlorine trifluoride.
NUT LIMITING
PIN
NICKEL PLUNGER
CKEL
STIRRER
SUPPORT
SHAFT HOUSING
MIDGET HOKE BELLOWS V A L V E SHAFT
N I C K E L SAMPLE POT
STIRRER THERMOCOUPLE
Fig. 1.-Freezing p,l/4' MONEL
WELL
point tube and stirrer. BELLOWS
1M' MONEL BELLOWS VALVES
CIFJ EXHAUST, 114'
SAE FLARE
SOOIUY
UB'SAE
FLARE
118' NICKEL
FLUSH
FLUOAIDE,
118. PELLETS..
TUBE.
TARE WEIGHT
SODIUM
114'
TUBIN0
G I ,JOG
FLUORIDE
TRAP.
A
MODIFIED
SAMPLE
Vol. 61
TUBE,
B Fig. 2.
boiler made it possible to shift the boiler composition iu convenient increments by this method. At the conclusion of a series of runs a t one total pressure, a fresh charge of pure chlorine trifluoride was used and the procedure re eated. The highly corrosive nature of chlorine triRuorid)e-hydrogen fluoride mixtures made it imperative that the time of contact of the mixture with the few silver soldered joints in the equilibrium still be minimized. At the conclusion of a run, the entire charge was transferred vapor-wise from the still to a welded nickel storage can cooled in liquid nitrogen. Even with this precaution corrosion was particularly severe when the mixture contained more than 50 formula yohydrogen fluoride. Several samples contained crystalline flakes identified as silver difluoride. Because of their low solubility and the analytical technique used, the presence of these salts did not affect the final results. Analysis.-The analysis of the equilibrium samples was based on the referential sorption of hydrogen fluoride on sodium fluor&. To serve as the sorption tower, a light trap of thin nickel tubing (Fig. 2A) was filled with 1/8 in. pellets of reagent grade sodium fluoride, evacuated and baked a t 400', pretreated with chlorine trifluoride, d?! filled with helium. The procedure consisted of vaporizing the sample mixture from the special sample tubes (Fig. 2B) through the tared sodium fluoride trap, absorbing hydrogen fluoride and venting chlorine trifluoride to a hood. Samples rich in hydrogen fluoride re uired warming of the sample tube in a beaker of water. Ipproximately 30 minutes was required to volatilize an average I-ml. liquid sample; a longer time was allowed for samples rich in hydrogen fluoride in order to avoid heating of the sodium fluoride trap. When
Results The data obtained from the warming and cooling curves are plotted in Fig. 3. All transition temperatures were determined under the orthobaric pressures of the mixtures. Chlorine trifluoride and hydrogen fluoride form a simple system with a single eutectic occurring a t -88.8' and 92.5 formula % HF. The expected exhibition of enantiotropy6by Below chlorine trifluoride was found at -83.1'. about 35 formula % HF the high temperature form of chlorine trifluoride separates on cooling the solution, while between 35 and 92.5 formula % HF the low temperature modification crystallizes. The temperature-composition data for the liquid-vapor equilibria obtained a t total pressures of 1.5, 2.0, 2.5, 3.0 and 4.0 atmospheres absolute are summarized in Table I and plotted in Fig. 4. Plotting the formula fraction chlorine trifluoride in the vapor against the formula fraction of chlorine trifluoride in the liquid shows that, within the experimental error, the separation factor at a given composition is independent of the total pressure in this range. Figure 4 illustrates the difficulty experienced in accurately assigning the azeotropic composition from the equilibrium curves alone; both liquid and vapor curves are flat and approach each other a t a low angle. Consequently, the azeotropic composition was determined by preparing a synthetic still charge mixture approximating the best estimate from the curves. From the degree of agreement between the liquid and vapor analyses, the composition could be more accurately assigned. On the basis of these measurements, the azeotropic compositions and boiling points have been assigned as in Table 11. Discussion All data have been presented on a formula per cent. basis because the actual molar concentrations are unknown. This complication arises because not only are both hydrogen fluoride and chlorine trifluoride associated but Penisler and Smith' report the formation of a weak molecular complex, ClF,.HF. I n the case of the solid-liquid and solidsolid phase equilibria, opportunity for theoretical treatment is particularly poor since association data are required in the temperature range of -75 to -90". For hydrogen fluoride, the necessary data for calculation of the degree of association have been obtained between 0 and 100" by Jarry and (6) J. W. Grisard, H. A . Bernhardt and G. D. Oliver, J . Am. Chem. Soc., 73,5725 (1951). (7) J. P. Pemsler and D. F. Smith, J. Chem. P h y s . , 32, 1834 (1954).
PHASE EQUILIBRIA OF THE SYSTEM CIFB-HF
August, 1957
-75
TABLE I LIQUID-VAPOR EQUILIBRIA OF THE SYSTEM CELORIDE TRIFLUORIDE-HYDROGEN FLUORIDE Total pressure, mm. abs.
Temp., OC.
1159 1151 1149 1148 1150 1149 1149 1150 1150 1150 1150 1149 1151 1151 1150 1150 1150 1530 1531 1530 1530 1530 1531 1529 1530 1530 1531 1530 1530 1531 1530 1530 1531 1531 1531 1899 1901 1901 1901 1901 1900 1900 1900 1900 2283 2281 2279 2277 2280 2287 2280 2280 2280 3039 3040 3041 3040 3041 3040 3040 3040
21.1 21 .o 21.0 20.4 20.0 20.0 20.7 21.6 24.7 29.1 29.3 27.7 24.2 24.3 24.4 24.5 23.4 28.6 28.7 29.4 29.4 29.8 30.4 30.9 31.4 31.9 32.8 33.6 34.5 36.4 28.0 27.8 27.8 27.9 28.7 33.7 33.7 33.3 33.4 34.1 35.6 37.7 42.3 34.4 40.7 39.4 38.8 38.9 38.7 38.7 39.4 41.8 47.0 47.5 48.4 47.1 47.2 48.7 49.6 51.0 56.8
1103
Equil. compositions, formula % Liquid
83.9 85.1 84.0 76.0 70.5 60.5 36.2 24.7 9.3 0.7 1.8 3.3 10.0 10.4 9.5 9.1 12.6 86.7 85.7 24.9 23.5 20.3 17.1 14.6 13.3
.. ..
6.6 5.9 3.0 77.8 72.0 65.0 52.8 33.6 81 .o
83.8 71.6 65.5 44.0 25.2 14.9 5.0 89.9 87.5 86.0 74.4 50.6 70.2 58.7 39.2 18.8 7.0 79.6 87.0 70.2 55.1
..
30.0 23.1 9.0
ClF,
Vapor
81.0 81.o 80.9 74.9 69.5 61.9 47.3 34.5 17.2 4.1 4.5 7.6 20.7 19.2 18.4
Y
.-80
a Q.
SOLID H F + SOLUTl
-85
Y
I
-90 -9 5 0
..
*.
23.1 20.6 17.2 15.0 12.9
60
80
70
FORMULA
90
100 H F
IHF.
60
58 56
54 52 50 48 46
b;w. 44 %
42
4
E
40
E
38 36
34
..
32 30
..
50
62
74.9 70.2 66.3 57.1 42.5 80.9 79.3 71.0 66.2 50.G 35.9 27.0 11.2 86.5 82.1 73.3 57.2 74.3 63.9 50.4 30.5 10.7 77.8 83.6 68.4 59.8 49.0 42.7 37.3 19.6
40
Fig. 3.-Solid-liquid equilibrium of the system chlorine trifluoride-hydrogen fluoride.
22.2 82.9 82.6 38.1 27.5
30
COMPOSITION,
..
..
20
10
28 26
"0
-
c\ /
01
i
0 2 03
04' 0 5 0 6 0 7
0 8 0 9 1.0
FORMULA FRACTION CIFa.
Wig. 4.--Liquid-vapor equilibria of the syfitem ClE'rHF. TABLE 11 AZEOTROPECOMPOSITIONS IN THE SYsiwi CHLORINE TRIFLUORIDE-HYDROGEN FLUORIDE Total pressure, mm. abs.
I150 1530 1900
2280 3040
B.p., OC.
19.9 27.7 33.3 38.5 47.1
Azeotrope composition, formula % ClE'
66.5 67.5 67.2 67.6 67.5
Davis,* but the formal representation of the temperature dependence is on an empirical basis and an 80" extrapolation of these data becomes somewhat questionable. The case for chlorine trifluoride is even less favorable since data of unknown accuracy covering a 15' temperature range would have to be extrapolated 100".
R. M. MCGILL,W. S. WENDOLKOWSKI AND E. J. BARBER
1104 e.10
2.00
I
t
1
1
I
l
I
1
HYDROGEN FLUORIDE-
I90
w
1
l
di
CHLORINE TRIFLUORIDE 160
::y, , , , ;hi 090 0
01
02
0.3
04
05
06
07
08
09
10
COMPOSITION. FORMULA FRACTION C I F s
Fig. 5.-Activity
coefficients in the system ClF3-HF.
The shape of the solidus curve in Fig. 3 clearly indicates a lack of compound formation and suggests instead the probable presence of a submerged miscibility gap and the accompanying large positive deviations from Raoult's law. Failure to detect the presence of the complex ClF3.HF reported by Pemsler and Smith' indicates that in this temperature range the bonding between chlorine trifluoride and hydrogen fluoride is weaker than the bonding between the hydrogen fluoride monomer units in hydrogen fluoride polymers or, expressed in terms of concentration, the concentration of hydrogen fluoride monomer in equilibrium with the ClF3.HF complex is larger than the concentration of hydrogen fluoride monomer in equilibrium with the hydrogen fluoride polymers. I n agreement with predictions of the general nature of the solution based on the solid-liquid equilibrium measurements, a minimum boiling azeotrope having a composition close to 67 formula % CIF3 for a wide range of pressures was found. The azeotrope can be separated into its components by passing the vapor through a bed of sodium fluoride which absorbs only the hydrogen fluoride. Activity coefficients of limited utility may be calculated for both chlorine trifluoride and hydrogen fluoride froin the liquid-vapor equilibrium data using the equation "/I
= YIP/XIPlO
(1)
where y1 is the activity coefficient of component 1, Y1 and Xi are the mole fractions of component 1 in the vapor and liquid, respectively, P is the total pressure, and PIois the vapor pressure of pure compoiient 1 a t the temperature of the measurement. This method defines the activity coefficient as the ratio of the observed vapor composition t o that of a vapor composition based on ideal behavior of the components in the liquid phase. Such a treatment assumes that the total pressure of the system can be represented by Dalton's law, ie., the sum of the partial pressures of the various molecular species present. I n the present system these species include not only chlorine trifluoride and hydrogen fluoride but also the polymers of hydrogen fluoride and the compound CIF3.HF. The analytical method used in this investigation does not distin-
Vol. 61
guish hydrogen fluoride from its polymers nor from hydrogen fluoride contained in C1F3.HF, but gives only the total formula composition of the sample mixtures in terms of hydrogen fluoride and chlorine trifluoride. Consequently, activity coefficients calculated from these data using equation 1must be interpreted in the sense that they are calculated, Le., on the assumption that only two molecular species are present. These coefficients for the 2.0 atmosphere data are shown graphically in Fig. 5 . There appears to be no unusual behavior a t the azeotrope composition. By use of the data in Table I and the data of Grisard, . Bernhardt and Oliver6 and Jarry and Davis,! the effect of temperature on the vapor pressure of solutions of fixed composition may be obtained. Plots of loglo P versus 1/T yield nearly straight lines. Considering the deviations to be within experimental error, the vapor pressure data have been expressed in Table I11 in terms of the equation loglo P = A - B/T (2) Besides the values of A and B , the apparent heats of vaporization, BR In 10, where R is the gas con-
C
TABLE I11 VAPORPRESSUREEQUATIONCONSTANTS FOR CLF~-HF SOLUTIONS Soln. comporn-
tion, formula % ClFa 0 10
20 30 40 50 07 80 90 95
100
BR In
Exptl.
temp.
range, O C. 32-63 24-56 22-52 21-50 20-49 20-48 20-47 21-48 21-49 22-50 22-51
A5 7.620 7.433 7.689 7.863 7.894 8.013 8.066 8.104 8.027 7.946 7.833
,
B, OK. 1393 1299 1366 1413 1420 1453 1468 1482 1462 1438 1407
10
ca1.7
moleb 6370 5940 6250 6460 6490 6640 0710 0780 6690 6580 6430
Calod.
heat of VaPn.
AK,.
Gal./ moleb 6370 6380 6380 6390 6390 6400 6410 6420 6420 6430
0430
"Ft mixing
cal.7
moleb 0 $440 +130 70 -100 -240 -300 -260 -260 -160 0
-
Pressure i i i inm. b The term "mole" is used wiLh r e w v:itions; the heat given is that for vaporization of sufficien1 solution to produce oiie GMV of saturated vapor assuming 1 GMV = 22.4litersatSTP. AssumingAH", = XHFAH"HF f XCIF~AH'CIF~ where AH"BF and L I H ~ Care I Fthe ~ values of B R In 10 for solutions having the compositions 100 formula % ' HF and 100 formula % ClF3, respectively. Assuming no change in volume on mixing, i.e., AH% = AHVx - B. R In 10. a
stant, have been listed for comparison with values of the heats of vaporization calculated using the equation AHV, = XEFAH'HF
+ XC~F~AH"CIF~(3)
where AHvXis the heat required t o vaporize from an infinite quantity of solution of composition XHFa t contant temperature and pressure, a quantity of solution which will give one gram molecular volume (22.4 liters a t STP) of saturated vapor; A H V ~and p A H V ~are ~ the ~ a values of BR In 10 for solutions containing 0.00 and 1.00 formula fraction ClFI, reare the formula spectively; and XHFand XCIF~ fractions of H F and ClF3, respectively. Implicit in the latter treatment are the assumptions (1) that the degrees of association of (HP)i and (ClF8)j do not change in going from liquid soh-
.
August, 1957
THERMODYNAMIC PROPERTIES OF SUBSTITUTED NAPHTHALENES
tion t o saturated vapor in any way which differs from the changes undergone by these materials in going from the pure liquids t o the pure saturated vapors, (2) that the actual effective number of individual molecules in the solution is equal t o the sum of the effective number of molecules in the pure liquids before they were mixed, (3) that the effective number of molecules in the gaseous mixture is the same as if the gases were unmixed, and (4) that the gas phase obeys the ideal gas law. Subject to these assumptions and the additional assumption of no change in volume on mixing, the seventh column in Table I11 gives the apparent heat of mixing of sufficient quantities of pure components to give one total "mole" of solution. The changes in the apparent heat of mixing with composition may be explained in terms of two major effects with opposite sign. The dissociation of (HF)i into hydrogen fluoride monomer and lower polymers requires the absorption of heat while the association of hydrogen fluoride monomer and chlorine trifluoride to form the weak complex, ClF4*HF, reported by Pemsler and Smith,s liberates 3.9 kcal. per mole of complex produced. Since significant quantities of the complex can only be produced if the partial pressure of hydrogen fluoride monomer is relatively high, dissociation of (HF)i induced by
1105
the addition of chlorine trifluoride to hydrogen fluoride rich solutions would be more pronounced than the formation of the complex, whereas the formation of the complex is more important in the chlorine trifluoride rich solutions. If, as is reasonable, the change in volume of the liquids on mixing is small, then the variations in the.apparent heat of mixing with composition are in the same sense as would be predicted. The shape of the activity coefficient curves in Fig. 5 tend to lend further credence to this explanation. The rapid initial increase in the activity coefficient of hydrogen fluoride with the addition of a small amount of chlorine trifluoride results from depolymerization of hydrogen fluoride polymer, and the subsequent leveling out of the activity coefficient of hydrogen fluoride a t 0.25 formula fraction chlorine trifluoride results from increased importance of the formation of the complex between chlorine trifluoride and hydrogen fluoride. It should be noted that there is no way of knowing the quantity of solution which constitutes one total "mole" without further experimental measurements. Acknowledgment.-The authors wish to thank J. W. Grisard and G. D. Oliver for performing the thermal analyses of the pure chlorine trifluoride and hydrogen fluoride.
THE LOW-TEMPERATURE THERMODYNAMIC PROPERTIES OF NAPHTHALENE, I-METHYLNAPHTHALENE, %METHYLNAPHTHALENE, 1,2,3,4-TETRAHYDRONAPHTHALENE, t~ans-DECAHYDRONAPHTHALENE AND cis-DECAHYDRONAPHTHALENE BY J. P. MCCULLOUGH, H. L. FINKE, J. F. MESSERLY, S. S. TODD, T. C. KINCHELOE AND GUYWADDINGTON Contribution No. 63from the Thermodynamics Laboratory, Pelrolum Experiment Station, Bureau of Mines, U.S. Department of the Interior, Bartlesville, Oklahoma Received April $8, 1967
As part of a program to provide basic thermodynamic information for important petroleum constituents, low temperature calorimetric investigations were made of the following six bicyclic hydrocarbons: naphthalene, I-methylnaphthalene, 2methylnaphthalene, 1,2,3,4-tetrahydronaphthalene,trans-decahydronaphthalene and cis-decahydronaphthalene. For each compound, measurements were made of the heat capacity in the solid and liquid states between 12 and 370"K.,the heat of fusion, the triple-point temperature, the cryoscopic constants, and the purity of the sample. Isothermal transitions that occur in 1-methylnaphthalene, 2-methylnaphthalene and cis-decahydronaphthalene were studied, and values of the heats and temperatures of transition were determined. Unusual effects of impurity on the thermal pro erties of l-methylnaphthalene and cis-decahydronaphthalene in the remelting" and melting regions were observed. #he absence of a reorted thermal anomaly in liquid cis-decahydrona'khalene was demonstrated. From the low temperature thermal data for each compound, values of the following thermognamic functions in the solid and liquid states were computed at selected temperatures from 10 to 370°K.: (Faatd - H0o)/T, (Hsatd - H"o)/T, Hsatd H'o, Saaaand C6atd.
-
An important fraction of the compounds found in the 180 to 230" boiling range of petroleum is composed of bicyclic hydrocarbons. Naphthalene and its derivatives probably comprise the bulk of this group of compounds. To provide basic information needed in the calculation of thermodynamic properties of bicyclic hydrocarbons, low temperature calorimetric studies were made of the following substances: naphthalene (I),1-methylnaphthalene (11), 2-methylnaphthalene (111), 1,2,3,4-tetrahy(1) F. D. Roesini, B. J. Mair and A. J. Streiff, "Hydrocarbona from Petroleum," ACS Monograph No. 121, Reinhold Publ. Corp., New,
York, N. Y., 1953.
CHs
03 03 03-cHa I
I1
0{111;CH2
IV
I CH2
I11
/CH2\/CHz\ CH2 CH
CH2
(!!HZ \CH/\CH/ (!XI
(!!€I2
V (trans) V I (cis)
