VAILABLE experimental information relating to the thermodynamic behavior of isobutane is meager. Dana and co-workers ( 4 ) reported the specific heat of the saturated liquid and the latent heat of I C . I ' I vaporization from 0" to 70" F. They also measured the density of the saturated liquid and gas and the vapor pressures from approximately 0" to 130" F. Seibert and Burrell (16) measured the vapor pressure of isobutane from 100"F. to the critical state. In connection with this work they also reported data concerning the critical pressure and temperature of this material. Burrell and Robertson (1) determined the vapor pressure a t temperatures in the vicinity of 0" F. Oudinoff (7) measured the atmospheric density of isobutane at 32" F. Parks et al. (8) reported heat capacity data for the liquid a t temperatures below 0" F. Daniel and Pierron (6)determined the ratio of the isochoric and isoB. H.SAGE AND W.N. LACEY baric heat, capacities a t atmospheric California Institute of Technology, Pasadena, Calif, pressure and approximately 60" F. This experimental information is insufficient t o establish the thermodynamic behavior of isobutane in the single-phase regions at many of the temperatures 250" F. The vapor pressure was also measured throughout and pressures commonly encountered in practice. The presthis temperature interval. From these primary data the ent investigation was made to obtain the data required variation in enthalpy (also known as heat content and as total to establish the thermodynamic behavior of this material heat), entropy, and fugacity with temperature and pressure under the conditions of temperature and pressure commonly have been calculated and are recorded in tabular form. encountered in the production of oil and gas from natural petroleum reservoirs. In an earlier paper of this series ( 1 4 , data concerning the isoProcedure baric heat capacity of isobutane a t atmospheric pressure were reported for temperatures from 70" to 320" F. I n the presMATERIALS.The isobutane used in this investigation was ent paper information resulting from the direct measurement obtained from the Philgas Division of the Phillips Petroleum of the isochoric heat capacity of isobutane in the two-phase Company. The results of their special analysis upon this region a t temperatures from 70" to 220" F. is recorded as well sample indicated that it contained less than 0.03 per cent imas results of direct measurement of the latent heat of vaporipurities. To avoid traces of noncondensable gas, the sample zation, a t temperatures as high as 170" F. The Joulewas repeatedly condensed as solid a t liquid air temperatures, Thomson coefficient of gaseous isobutane was measured a t under a pressure only slightly greater than its sublimation temperatures from 70" to 250" F. The specific volume in the pressure a t the temperature involved. The material after single-phase regions was determined at pressures from atthis treatment exhibited approximately 0.2 pound per square mospheric to 3000 pounds per square inch and from 70" to inch change in vapor pressure from dew to bubble point. A sample of this material was further purified by fractionation in a vacuum-jacketed column packed with glass rings (16). The middle fraction from this distillation still showed a variation of approximately 0.15 pound per square inch change in Specific volumes, vapor pressures, heat vapor pressure from dew point to bubble point a t 160" F. capacities, Joule-Thomson coefficients, METHODS.The methods employed for the pressure-voland latent heats of vaporization were ume-temperature measurements were described recently ( l a ) . It is believed that the pressures which were measured by use experimentally determined at temperaof a calibrated fluid-pressure balance were determined within tures from 70"to 250"F. From these data, an uncertainty of 0.2 pound per square inch. The temperatogether with some other pertinent pubtures were ascertained by means of a copper-constantan therlished information, the changes in enmocouple used in conjunction with a potentiometer having a thalpy, entropy, and fugacity for changes sensitivity of 0.1 microvolt. It is believed that the temperatures reported are known within 0.1' F. in relation to the inin temperature and pressure were calternational platinum scale. However, a t the highest temculated. The results of the experimental peratures reported here, a somewhat larger deviation may work, as well as values of the derived have occurred. The volumes of the samples were determined properties mentioned above, are presented with an Uncertainty of 0.1 per cent. A comparable error was in tabulated and graphical form.
A
Phase Equilibrium in Hydrocarbon
Systems
Thermodynamic Properties of Isobutane'
1 This is the twenty-first paper in this series. Previous articles appeared during 1934 to 1937, incluaive.
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P I
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- 15.0
50
100 TEMPERATURE
FIGURE3.
TEMPERATURE
' F.
FIGURE2. ISOBARICTHERMAL EXPANSION OF CONDENSED LIQUID
involved in the measurement of the weight of sample employed in studying the behavior of the superheated gas. The error in weighing the samples used in the study of the liquid region was less than 0.03 per cent. The apparatus employed for the measurement of the isochoric heat capacity in the two-phase region was described previously (IO). Briefly, the method consists of measuring the temperature rise resulting from the addition of a known amount of electrical energy to a steel container of low heat capacity nearly filled with liquid isobutane. Heat losses from the container were avoided by the use of an adiabatic vacuum jacket. It is believed that the heat capacity data reported for the saturated liquid are not in error by more than
I
200 -
150 F.
RESIDUALVAPORPRESSURE-TEMPERATUR DIAGRAW
1 per cent, except a t temperatures in excess of 200' F., where the uncertainty may be somewhat larger. A modification of this calorimeter was employed in the measurement of the latent heat of vaporization. I n principle, the method was similar to that employed by Osborne and Van Dusen (6) in their study of the latent heat of vaporization of ammonia. I n the present case the superheat of the vapor evolved from the calorimeter was smaller than that encountered by Osborne and Van Dusen. I n general, all of the corrections for variation of the conditions within the calorimeter from the ideal isobaric isothermal vaporization amounted to less than 2 per cent for all of the values obtained except those a t temperatures in excess of 170" F. These latter values were not employed in calculating the thermodynamic properties tabulated here. It is believed that the latent heat of vaporization was determined with an uncertainty of less than 1.0 per cent throughout the temperature range below 170" F. The equipment used in measuring the Joule-Thomson coefficient of gaseous isobutane was also described in another paper (9). I n the present case a pressure change of approximately 1.5 pounds per square inch was maintained across a porous radial-flow plug, and the resulting temperature change was measured by means of a multijunction thermocouple. Care was taken to ensure adiabatic flow and to avoid appreciable changes in velocity. The small pressure changes employed permitted measurements to be made reasonably close to the state of saturated gas. However, saturation conditions could not be completeIy realized because traces of oil in the apparatus caused the appearance of a liquid phase a t pressures slightly below vapor pressure. It is believed that the Joule-Thomson measurements were made with an uncertainty of less than 2 per cent except in the vicinity of saturated gas.
Results The results of the pressure-volume-temperature measurements upon the condensed liquid are depicted iil Figure 1. The experimental points shown were taken at random from
JUNE, 1938
INDUSTRIAL A N D ENGINEERING CHEMISTRY
675
perature relation of isobutane can be approximately described by the following expression: hgio PS = 5.47567
- 2024.99 T
If this equation is used as a reference basis, the residual vapor pressure may be evaluated from the experimental measurements and Equation 1 by the following relation: P=PIL-P
I
(2)
Figure 3 shows the residual vapor pressure as a function of temperature. This method of presentation allows a more accurate representation than is possible by direct plotting of the relation of the vapor pressure to temperature. The full line represents the values chosen for the vapor pressures recorded in Tables I1 and 111. Although the authors' points are consistent with one another within 0.05 pound per square inch, it
FIGURE 4. COMPRESSIBILITY FACTOR FOR GASEOUS ISOBUTANE
three independent sets of measurements upon samples of isobutane of markedly different weight. The data obtained TEMPERATURE 'E from these different samples agreed with one another within FIGURE 5 . RESIDUAL SPECIFIC VOLUME OF 1 part in 2000 throughout the range of pressures and temperaGASEOUSISOBUTANE AT ATMOSPHERIC tures indicated in Figure 1. It is believed that the specific PRESSURE volumes of the condensed and saturated liquid which are recorded in parts of Tables I1 and 111,within the range of temis believed that an absolute uncertainty as great as 0.5 pound perature covered by the experimental work, are known with per square inch may exist in the values at the higher temperaan accuracy of 0.1 per cent. The value for specific volume of tures. The measurements of Dana et al. (4) agree with the the saturated liquid a t 70" F. is 0.2 per cent lower than that measurements of the authors with an average deviation of found by Coffin and Maass (3) and 0.35 per cent lower than that obtained by Dana et aE. (4). For some purposes the thermal expansion of liquid isobutane is of i n t e r e s t . Figure 2 presents the variation in the isobaric thermal expansion with temperature for a series of p r e s s u r e s . These data were obtained by residual graphical differentiation of the specific volume data recorded in a part of Table 11. This d i f f e r e n t i a t i o n was carried out with sufficient precision to avoid uncertainty in the isobaric thermal expansion beyond that involved in reading values from Figure 2. These data indicate a rapid increase in the thermal expansion with an increase in temperature, and a decrease in the value of this derivative with an increase in pressure. The latter effect becomes more pronounced a t the higher temperatures. This behavior is similar to that found for other paraffin hydrocarbons. The vapor pressures were determined from a special set of measurements in which the sample was maintained a t a quality of approximately 0.1 throughout the temperature range. There was a change of approximately 0.15 pound per square inch in vapor pressure within the course of an isothermal condensation, but the change in pressure was almost PRESSURE LB. PER %IN. entirely limited to the states in the vicinity FIGURE 6. JOULE-THOMSOX of saturated gas. The vapor pressure-temCOEFFICIENTS FOR GASEOUS ISOBUTANE
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perature measurements and the Joule-Thomson data. The results reported by Dana et al. (4) upon the 0.80 d e n s i t y of t h e saturated gas, as evaluated from the heat of vaporization measurements, are included for comparison. The agreement is within 055 the experimental uncertainty of their measurements. Oudinoff (7) measured the atmosc 0 50 pheric density of isobutane, thus permitting another check upon the accuracy of the volumetric data de0.45 rived from the Joule-Thomson measurements. The curve of Figure 5 represents the variation in the residual 040 volume with temperature at atmospheric pressure as based on the JouleT h o m s o n d a t a a n d t h e pressurevolume-temperature measurements a t 220" F. This indicates that there is TEMPERATURE 'E a satisfactory agreement (0.05 per cent) between Oudinoff's value and FIGURE 7. HEATCAPACITY OF GASEOUS ISOBUTANE the present experimental measurements. The experimental data for the JouleThomson coefficient of gaseous isobutane are presented in Figure 6. The agreement of the experimental points with the smooth curves is not as good as might be desired, but is still comparable with the experimental precision of the measurements (2 per cent). Figure 6 indicates that the JouleThomson coefficient is finite for this gas at infinite dilution. A knowledge of the Joule-Thomson coefficient as a function of state and of the isobaric heat capacity as a function of temperature at a single pressure is sufficient to ascertain completely the heat capacity throughout the range of pressure for which the Joule-Thomson data are available. The method employed (9) in such calculations is based uDon the progressive graphical inteapproximately 1.1 pound per square inch. The measuregration of the following general thermodynamic relation : ments of Seibert and Burrell (16) are included for comparison. They are from 7 to 14 pounds per square inch higher than the (3) values reported in this paper. As in the case of earlier work upon n-butane (13),difficulty Data upon the isobaric heat capacity of isobutane a t atmoswas encountered in determining directly the pressure-volumepheric pressure were recently published (14). These values, temperature relations of the superheated gas a t the lower together with the Joule-Thomson data presented in Figure 6, temperatures. The direct measurements below 190" F. were were used to calculate the heat capacity for the superheated disregarded in favor of an indirect evaluation by the use of region, as recorded in graphical form in Figure 7 . It is beJoule-Thomson coefficients, although the maximum disagreelieved that these heat capacities are known with an uncerment between the two sets of values was only 0.6 per cent, extainty of approximately 1per cent a t the lower pressures, but cept in the immediate vicinity of saturated gas where diverthe uncertainty is larger a t states in the vicinity of saturated gences as large as 1.0 per cent were encountered. The regas. sults of the direct measurements a t temperatures of 190" F. Figure 8 presents the experimental results obtained for the and higher are shown by the experimental points in Figure 4. isochoric heat capacity in the two-phase region a t a quality The measurements for 229" F. were used as a basis in calcuof approximately 0.1. These isochoric data were converted lating the values obtained from the Joule-Thomson coeffito isobaric heat capacities for the saturated liquid by means cients. These values are represented by the curves for the of the following 677) other temperatures. The methods employed in such calculations have been described recently (13). The agreement between the direct measurements and the behavior calculated in this manner are considered as partial indication of the thermodynamic consistency of the direct pressure-volume-tem-
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It is believed that these heat capacity data are trustworthy within 1 per cent t h r o u g h o u t the temperature range recorded here. The measurements of Dana et al. upon the heat capacity of isobutane liquid under saturation conditions are included for comparison. These data, which were reported as heat capacities under saturation conditions, were converted to isobaric heat capacities by the equation:
The ageement (approximately 8 per cent) is not so good as would be expected. The authors are unable to account for the large value of (dCb/dT) found by Dana et al. (4) for isobutane as c o m p a r e d to the much smaller value the same authors found for propane and n-butane. T h e p r e s e n t heat capacitv " measurements are in good a g r e e m e n t with v a l u e s ex-
TEMPERATURE
'6
ENTHALPY-PRESSURE COEFFICIENT FOR GASEOUS ISOBUTANE FIQURE 9. ISOTHERMAL
2 This expression may be derived from general thermodynamic relations 81 follows: For any equilibrium isothermal change in state,
(2) ~
av
T ( %aTp )
v
T
In the two-phme region for a pure substance where (alp/ bTP)v is independent of volume and equal to (d2PU/dT2). the above expression may be integrated for a change in state from Y to Yb:
In a recent publication ( 2 1 ) it was shown that
A combination of Equations b and c with the following general thermodynamic relation yields Equation 4:
F.
TEMPERATURE
FIGURE 10. HEATOF VAPORIZATION TABLEI. JOULE-THOMSON COEFFICIENTS AND ISOBARIC HEATCAPACITIES FOR ISOBTJTANE GAS Abs. Pressure -70' F . 7 --100' F.-130O F.Lb./sq. i n . P CP Ir CP 9 CP Satd.gas O.341Sa 0.4290b 0.3051 0.4540 0.2740 0.4804 Satd. liquid 0.5631 .... 0.5972 0.6353 0 0.1500 0.3848 0.1310 0.3946 0.1119 0.4058 14.7 0.2088 0.3896 0.1692 0.3995 0.1384 0.4096 20 0.2308 0.3941 0.1817 0.4011 0.1478 0.4113 40 0.3183 0.4192 0.2307 0.4143 0.1802 0.4196 60 .... . . . . 0.2785 0.4363 0.2105 0.4313 80 .... 0.2380 0.4470 100 .... .... 0.2645 0.4692 125 .... 150 .... 175 .... .... 200 .... .... .... 225 .... .... .... .... .... 250 300 .... .... .... .... .... 350 .... .... .... ..,. .... 400 .... .... 0 Joule-lhomson coefficients are ex ressed a s F./ ound per square inch. b Heat cat acities are expressed as t. U. per poun@ F. C Joule-Tfomson coefficients for 250° F. were obtained by extrapolation.
....
.... .... ..,.
....
.... .... ....
....
....
....
.... .... .... ....
....
....
.... .... .... .... ....
F.-190° CP P 0.2495 0.5112 0.2329 0.6820 0.0928 0.4170 0.0739 0.1141 0.4212 0.0928 0.1217 0.4227 0.099; 0.1478 0.4292 0.1220 0.1724 0.4371 0.1435 0.1930 0.4468 0.1610 0,2113 0.4585 0.1756 0.2288 0.4767 0.1912 0.2432 0.4997 0,2039 .... .... 0.2150 0.2239 0.2322
....
8.
.... .... .... .... .... .... .... ....
.... .... .... .... .... .... ....
.... ....
--160° P
....
.... ....
.... .... .... ....
....
.... ..,. ....
.... .... .... ....
F.-? -220° CP P 0.5475 0.2219 0.7413 .... 0.4293 0.0548 0.4332 0.0729 0.4347 0.0789 0.4406 0.1003 0.4471 0.1188 0.4542 0.1348 0.4622 0.1479 0.4735 0.1619 0.4869 0.1740 0,5019 0.1842 0.5200 0.1926 0.5431 0.2003
.... .... ....
0 2076 0.2189
....
F.CP
0.600 0.8193 0.4419 0.4458 0.4473 0.4529 0.4590 0.4654 0.4722 0.4813 0.4912 0.5020 0.5140 0.5275 0.5428 0.584
.... ....
--250' 9
0.216C
F.CP
....
0.035 0.4555 0.054 0.4594 0.0599 0.4609 0.0797 0.4665 0.0965 0.4725 0.1112 0.4787 0.1230 0.4853 0.1355 0.4938 0.1466 0.5028 0.1568 0.5123 0.1645 0.5222 0.1732 0.5326 0,1800 0.6437 0.1922 0.569 0.2028 0.602 0.212
....
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t r a p o l a t e d from measurements by Parks el al. (8) a t temperatures below 0 ' F. F o r 'convenience, values of the Joule-Thomson coefficient and the isobaric heat c* pacity for t h e TEMPERATURE superheated g a s OF HEATOF FIGURE11. COMPARISON have been recorded VAPORIZATION AS DETERMINED BY a t a n u m b e r of DIFFERENT METHODS temDeratures for a series of pressures in Table I. The isobaric heat capacity of the saturated liquid is also included. From these-data the isothermal change in enthalpy with pressure may be calculated from the general thermodynamic relation:
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ploying the volumes of the saturated liquid and saturated gas in conjunction with the rate of change of vapor pressure with temperature. The change in the latent heat with temperature may be calculated from heat capacity data. If a single value of latent heat is chosen as a reference basis, the latent heat a t other temperatures may then be determined from the following expression:
OF:
(g)
=
- PCP
A comparison of results obtained by these methods of evaluation with the values reported in Tables I1 and I11 has been made in Figure 11. I n this case the latent heat a t 70OF. was used as a reference basis for the calculation based upon
(6)
The results of such calculations are p r e s e n t e d in Figure 9 where the isothermal enthalpy-pressure coefficient is d e p i c t e d as a function of temperature for a series of pressures. These data indicate the characteristic minimum in the value in this coefficient for the saturated gas which was found for both propane and n-butane. A curve for this derivative for infinite dilution is included in Figure 9. The results of the direct experimental measurements of the heat of vaporization are presented in Figure 10. With the exception of the data in the v i c i n i t y of 170" F., v a l u e s show an average deviation of approximately 1 per cent. The measurements of Dana et ab. (4) are included for comparison. T h e a g r e e ment between the two sets of data is satisfactory, and the heat of vaporization a t 70' F. thus determined was used as the basis of the entropy and enthalpy of the superheated gas reported in Tables I1 and 111. The latent heat of vaporization may be evaluated from the data a v a i l a b l e other than by use of the direct experimental measurements. It may be calculated by means of the Clapeyron equation, em-
ENTROPY 0 05
0 10
0 I5
I 0.20
I 0.25
B. T. U. PER LB. PER "E 030
0.35
0 40
DIAGRAM FOR ISOBUTANE FIGUREI 12. TEMPERATURE-ENTROPY
I OA5
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Tables I1 and I11 were taken from the directly measured values. From 120" to 175O F., values derived from the specific heat data were employed. Above 190" F. the greatest weight was given to the Clapeyron equation values.
TABLE111. THERMODYNAMIC PROPERTIES OF ISOBUTANE, SATURATED GASAND SATURATED LIQUID -Saturated
P
t
V
40 50 60 70 80 90 100 125 150 175 200 225 250 275 300 325 350 375 400
63.0 76.5 88.1 98.2 107.3 115.5 123.8 139.8 154.2 167.0 178.3 188.7 198.3 207.3 215.6 223.5 231.0 238.1 244.9
2.210 1.7813 1.4904 1.2796 1.1198 0.9947 0.8949 0.7103 0.5864 0.4979 0.4305 0.3769 0.3327 0.2954 0.2633 0.2325 0.2110 0.1888 0.1686
GasH S 146.4 0.2803 151.11 0.2818 154.82 157.97 0.2831 0.2841 160.81 0.2852 163.33 0.2862 165.73 0.2871 170.44 0.2889 174.49 0.2906 178.03 0.2923 181.0 0.2938 183.8 0.2951 185.8 0.2957 187.3 0.2959 188.7 0.2959 189.6 189.6 0.2954 0.2941 189.5 0.2920 188.7 0.2897
f 37.16 45.83 62.52 54.28 70.56 78.42 86.11 104.73 122.60 139.82 156.45 172.53 188.19 203.4 218.1 232.5 246.7 260.7 274.5
-Saturated
V
0.02838 0.02888 0.02973 0.02932 0.03013 0.03049 0.03088 0.03167 0.03245 0.03331 0.03412 0.03496 0.03578 0.03663 0.03748 0.03838 0.03935 0.04036 0.04143
LiquidH S 1.64 0.0032 9.30 0.0173 21.96 16.01 0.02957 0.0403 27.34 0.0499 32.37 0.0586 37.57 0,0674 47.89 0.0844 57.36 0.0998 66.06 0.1136 73.94 0.1259 81.42 0.1373 88.51 0.1578 0.1478 95.26 101.7 0.1671 108.0 0,1760 114.1 0.1846 120.1 0.1928 126.1 0.2009
the specific heat data because of the excellent agreement between the direct measurements of Dana et al. and the authors' direct measurements a t this temperature. Above 130" F. the direct measurements are slightly larger than the values chosen. This discrepancy may be due in part to the larger heat losses attending the calorimetric measurements and in part to the increasing importance of the volumetric correction terms which must be applied to the experimental data.
%-S
B.T. U. PER LB PER
FOR
Thermodynamic Calculations
From the primary data discussed in the preceding section, the variations in enthalpy, entropy, and fugacity with pressure and temperature were calculated. These data are more than sufficient to establish the thermodynamic behavior of this material. Their thermodynamic consistency with one another was discussed earlier and is within the absolute accuracy of measurement in all cases, except possibly for the volumetric behavior in the two-phase region below 100" I?. The values of enthalpy and entropy recorded in Table I1 are based upon an arbitrary value of zero for the enthalpy and entropy of the saturated liquid a t 60" F. The methods employed in these calculations were described previously (11, 13). Residual graphical methods were employed, and the data recorded in Tables I1 and I11 are consistent within 1part in 1000, except in the immediate vicinity of saturated gas where a somewhat larger degree of inconsistency is encountered. The degree of consistency was established by application of several different general thermodynamic relations for changes of state in and between the one- and twophase portions of the system. Numerous diagrams may be drawn from the tabulated data to represent the thermodynamic behavior of isobutane. Figure 12 is a temperature-entropy diagram for this material including the liquid, gas, and two-phase regions. Lines corresponding to constant values of pressure, specific volume, and enthalpy are shown, and lines of constant quality are included in the two-phase region. The diagram was extended to include the critical region, using the data of Seibert and Burrell (16) to establish the critical temperature. The critical volume was determined by an extrapolation of the average density (2) of the saturated liquid and gas to the critical temperature. The critical pressure was obtained by an extrapolation of the vapor pressure data recorded here, to the above mentioned temperature. Because of the rapid changes in the properties of the system in the single-phase regions in the vicinity of the critical state, no attempt was made to depict these regions on Figure 12. Figure 13 shows the isothermal change in enthalpy and entropy from saturation for a series of temperatures. This diagram is included to assist in comparison of the behavior of liquid isobutane with that of n-butane (13).
Nomenclature
F:
T CHANGES IN ENTHALPY AND ENTROPY FIGTJRE 13. ISOTHERMAL t CONDENSED LIQUIDS
P
p
'
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Below 130" F. the heat of vaporization calculated from the Clapeyron equation does not agree with the experimental measurements as well as might be desired. This lack of agreement (1.1 per cent a t 70" F.) is probably due to small errors in vapor pressure measurements and to some uncertainty as to the specific volume of the saturated gas. At higher temperatures the values from the Clapeyron equation are slightly larger than those chosen, but again the agreement becomes better a t temperatures above 190" F. The latent heat data a t the lower temperatures reported in
-Hvv
s
f
cp
cv
L p
= temp., =
temp.,
O
F. abs. ( " R.)
O F .
= pressure, lb./sq. in. abs. = residual pressure, lb./sq. in. specific volume, cu. ft./lb. residual volume, CU. ft./lb. = enthalpy, B. t. u./lb. = entroDv. B. t. u. Der lb./" F. abs. fugac'lcy, lb./sq. h. . = isobaric heat capacity, B. t. u. per lb./" F. = B. t. u. per lb./" F. = isochoric heat capacit = heat of vaporization, t. t / l b .
= =
Subscript b
=
g.
coefficient, F./lb. per sq. in. a system at bubble point (satd. liquid)
= Joule-Thomson
Subscript d = a system at dew point (satd. gas) Superscript ' = a system in a single-phase region Superscript " = a system in the two-phase region
JUNE, 1938
INDUSTRIAL AND ENGINEERING CHEMISTRY
Acknowledgment This investigation was made under the auspices of Research Project 37 of the American Petroleum Institute. Financial support and cooperation from this institute are acknowledged. W. R. Mendenhall carried out the measurements of the JouleThomson coefficients reported. T. Vermeulen determined the directly measured heats of vaporization and the heat capacities of the liquid. The assistance of Louise M. Reaney and D. C . Webster is acknowledged in connection with the calculations.
Literature Cited (1) Burrell and Robertson, U. S. Bur. Mines, Tech. Paper 214 (1916). (2) Cahetet and Mathias, Coompt. rend., 102, 1202 (1886).
681
(3) Coffin and Maass, J . Am. Chem. SOC.,50,1427 (1928). (4) Dana, Jenkins, Burdick, and Timm, Refrig. Eng., 12, 387 (1926). (6) Daniel and Pierron, Bull. SOC. chim., 21,801(1899). ( 6 ) Osborne and Van Dusen, U. S. Bur. Standards, Sci. Paper 315
(1918). (7) Oudinoff, Bull. S O C . chim. Belg., 23, 266 (1909). (8) Parks, Shomate, Kennedy, and Crawford, J . Chem. Phys., 5, 369 (1937). (9) Sage, Kennedy, and Lacey, IND. ENQ.CHEM.,28,601 (1936). (IO) Sage and Lacey, Ibid., 27, 1484 (1935). (11) Sage, Webster, and Lacey, Ibid., 26, 1213 (1934). (12) Ibid., 29, 658 (1937). (13) Ibid., 29, 1188 (1937). (14) Ibid., 29, 1309 (1937). (15) Seibert and Burrell, J . Am. Chem. SOC.,37, 2683 (1915). (16) Young and Jasaitis, Ibid., 58,377 (1936). RECEIVED December 27. 1937.
Nitroglycerin and Ethylene
Glycol Dinitrate Vapor Pressures of Binary Solutions
S
INCE ethylene glycol dinitrate is now extensively used
in commercial explosives in solution with nitroglycerin (glycerin trinitrate) t o increase freezing resistance, vapor pressure data on binary solutions of these compounds are of industrial importance, particularly from the standpoint of volatility. A search of the literature revealed that no measurements of the vapor pressures of such solutions have been reported. As for the individual esters, an examination of the literature showed that vapor pressures of nitroglycerin have been reported by Marshall (4) and Peace (8), Chiaraviglio and Corbino (Z), Naoum and Meyer (IO), and Crater (W), and the vapor pressures of ethylene glycol dinitrate by Crater (2) and Rinkenbach (13). The values reported by the various authors are not in agreement with one another. Marshall (5) gave a review and discussion of previous work and selected from a diagram of all the published results those values of the vapor pressures which he considered most probable. Several papers of a polemical nature have been published by Marshall (6, 6, 7) and by Naoum and Meyer (9, 11, 22). This work was accordingly undertaken to determine the vapor pressures of solutions of nitroglycerin and ethylene glycol dinitrate as well as to check the values reported in the literature for the vapor pressures of the pure compounds.
Preparation of Nitric Ester Solutions ANALYSISOF MATERIALS BEFORE NITRATION. Ethylene glycol and glycerin samples of high purity were used in the preparation of the nitric esters. The following table shows the analysis of these raw materials, carried out according t o International Standard Specification Methods:
J. D. BRANDNER Atlas Powder Company, Tamaqua, Pa.
The vapor pressures of nitroglycerin were measured from 30" to 50" C., and of ethylene glycol dinitrate from 20" to 50" C., using a dynamic method. By extrapolation of the straight lines obtained by plotting the data as logarithm of vapor pressure us. reciprocal of the absolute temperatures, final values of the vapor pressures of the pure niitric esters were selected from 10" to 50" C. Measurements at 40" and 50" C. of the concentration of vapors from binary solutions of nitroglycerin and ethylene glycol dinitrate indicated them to be ideal. Equations were accordingly derived using the Clapeyron-Clausius and Raoult equations for calculating the vapor pressure of any binary solution at any temperature from 10" to 50" C.
Glycerin
Specific gravity Silver teat Analysis, % by wt.: Char Ash Chlorides Acidity (50 cc. required), cc. 1 N NaOH
Ethylene Glycol 1.1176 Light red-brown No color, no ppt. color, no ppt. 1.2620
0,029 0,003 0.0002
0.001 0.0004 0.0001
0.04
0.01
PUREXITRIC ESTERS.The mixed acid used in the nitrations had the following composition: total sulfuric, 49.99 per cent; total nitric, 52.44 per cent. Samples of the esters used were from laboratory nitrations of two glycerin and ethylene glycol samples and were given special treatment in purification and drying. After having been washed with a solution of sodium carbonate for the purpose of neutralizing
