KOTES
March, 1961 solvent may be different from that in water.'O Determinations of the heat of neutralization of strong acids by strong bases in solutions of dioxane-water varying from 0 to 75 volume yo dioxane show a variation of only 2.4%. The condition of solvation of the proton and hydroxide ion must thus differ only by very little in the two solvents. Emerimental The apparatus and procedures for enthalpy titrations are essentially those described hy Jordan and Alleman." For each run, 100 ml. of 0.01 M sodiuni hydroxide in the given solvent was titrated with 1. O M hydrochloric acid in the same solvent. The results assenibkd in Table I are the avrrage of 5-9 determinations a t each mole fraction of dioxane.
TABLE I HEATSOF NEUTRALIZATIOI. I N WATER-DIOXAXE ---Dioxane--Vol C L
0 25 50 75
Mole fr
0 0 0 0
00 066 li4 388
4 H neut.
13 13 13 13
57 65 70 89
f0 f0 f0 f0
02 14 04 04
(10) Private conimunications. (11) J. Jordan and T. G. Alleman, 9naZ. Chem., 29, 9 (1957).
SOLUBILITY AND THERMODYNAMIC FUNCTIONS OF ETHYLENE IK DIETHYL SULFATE BY A. M. TRUCHARD, H. G. HARRIS ANJ) D. M. H~MMELBLALI Department of Chemical Engineering, The Unzverszty of Texas, A u s t t n
=
log
575 3.04666 - o.132845 5.9302
x
10-4
was calculated by the least squares technique to give an excellent fit (maximum deviation from 0 to 80" was 1.51Yo). With the aid of this equation, the partial molal heats of solution were calculated as
and the partial molal entropies of solution were calculated as
The calculated thermodynamic values are listed in Table I. Womenclature and discussion related to equat,ions2 and 3 is in reference 1. TABLE I HENRY'SLAWCONST.4NTG AND THERMODYNAMIC DATA FOR ETHYLENE - (EL - Bo), tal./ - Bo), %(""-_) - (RLO 1.
"C.
mole f r I n EtnSOh
cal./g. molea
In EtL304 In H20
( g . mole)('K.) b I n EtsSOa In H20
2430 5260 17.9 36.4 0 68.71 2300 4390 16.8 33.3 20 92.24 2270 4170 16.7 32.6 25 98.73' 30 105.1 2240 3880 16.6 31.7 2180 3340 16.4 29.9 40 117.3 2080 2100 16.1 26.0 60 144.2 1995 850 15.9 22.3 80 172.8 "infinite dilution." * Standard State: a Standard State: mole fraction = 1.0. Calculated.
12, Texas
Reeelired October 84, 1960
As part of a study of the reaction of olefins with sulfuric acid, a thorough investigation of the solubility of ethylene in diethyl sulfate has been made. From the solubility data the partial molal heats and entropies of solution have been calculated. (a) Apparatus.-The experimental apparatus consisted of two calibrated volumetric bulbs which contained approximately equal volumes of gas after 200 ml. of diethyl sulfate had been added to the larger bulb. The flasks were immersed in a thermostat (good to f0.05O) and attached to a mercury manometer. The system was evacuated by prolonged pumping, and ethylene at about 2 atm. pressure was added to the empty smaller flask. Pressures in the system were measured, and then the ethylene was admitted to the bulb coiitaining the diethyl sulfate. The system was allowed to reach equilibrium while being stirred. For each temperature successive additions of ethylene were brought into contact with the diethyl sulfate a t pressures ranging from 50 to 1300 mm. Plots of corrected partial pressure of ethylene us. mole fraction ethylene dissolved in diethyl su!fate could be fit with excellent precision by straight linrs. T h r diethyl sulfate wap Eastm:tn practical grade purified by washing with Xa,CO, solution and drying with CaC11. A t temperatures much higher than 100" diethyl sulfate begins to degrade so the maximum temperature reported is 80".
(bj Results.--Tahle I shows the Henry's law constants from 0 to 80" computed as X = p (atm.)/x (mole fraction). Since the logarithms of the Henry's law constants were not exactly a lirieur function of temperature, a curve to fit the values of x of the form
Discussion The reliability of t,he Henry's law constants was excellent since the standard deviation for X for any temperature was less than 0.01. The accuracy of the apparatus was tested by measuring the solubility of ethylene in water at 0'. X was determined t'o be 5.28 X lo3, which is 4.5yoless than the value of Winkler.2 Winkler's value may be high, however, since the data of Davis and McKetta3 and Bradbury, McKult8y,Savage and McS ~ e e n e y ,when ~ extrapolated to lower tempera tures, yields a lower value (less than 5 X lo3) for
x.
The partial molal heats and entropies of solution of ethylene in diet,hyl sulfate do not change as much mit,h temperature as do corresponding values for ethylene dissolved in water, which have been calculated from the data in reference 5, although the trend with temperature is the same. The values of (& - SG)in diethyl sulfate are onehalf as negative as those in water and correspond closely t'o the values shown in Table I1 for other non-polar compounds. The values in Table I1 have been computed from the data of Horiuti6 (1) D. hl. Hirnmelblau, J . Phya. Chem., 63, 1863 (1959). ( 2 ) L. W. Winkler, Z. phusik. Chem.. 55, 350 (1906). (3) J. E. Davis and J. J. McKetta, J . Chem. Eng. D z t a , 6 , 37-1 (1960). (4) E. J. Bradbury, D. RlcNulty, R . L. Savage and E. E. McSweeney, I n d . Eng. Chem., 44, 211 (1952). (5) D. M. Hirnmelblau and E. Arends, Chem. Iny. Tech., 31, 791 (1959).
576
COMMUNICATION TO THE EDITOR
An interpretation of the thermodynamic data is that the quasi-ice-like structures caused by the dissolution of ethylene in water are more rigid or penetrate to a greater distance from the ethylene molecule than those created in diethyl sulfate and other non-polar compounds. Presumably this is due to the hydrogen bonding in water. Furthermore, the type of structure in the non-polar compounds is quite stable as the temperature changes; the thermodynamic functions for the compounds listed in Table I1 have essentially constant values over the temperature range for which data are available. The arrangement of the structure surrounding the ethylene molecule in diethyl sulfate appea:rs to be independent of the nature of the solvent for both diethyl sulfate and those compounds listed in Table 11. (6) J. Horiuti, Bull. Insf. Phys. Chem. Research (Tokyo), 9, 697 (1930); Sei. Papers Inst. Phys. Chem. fiesearch (Tokyo), 17,125 (1931).
Vol. 65
TABLE I1 PARTIAL MOLALHEATSA N D ENTROPIES OF SOLUTION FOR ETHYLENE DISSOLVED IN NON-POLAR SOLVENTS~ (0 TO 40') - PLO- ZG), - (ISL - QG), caI./ Solvent
cA./g. moleb
(g.
mole)('K.)
c
2420 16.5 CCll CHaCOOCH3 2260 16.5 2220 16.2 CaHa (CHahCO 2130 16.1 Standard State: Calculated from data of reference 6. "infinite dilution." 0 Standard State: mole fraction = 1.o.
Acknowledgment.-This research was supported by a grant from the Petroleum Research Fund administered by the American Chemical Society. Grateful acknowledgment is hereby made to the donors of this fund.
COMMUNICATION TO THE EDITOR SPECIFIC REFRACTIVE INCREMENT OF POLYPROPYLENE I N a-CHLORONAPHTHBLEKE
to 125') or the value of -0.227 cc./g. at 125" reported by Kinsinger and Hughes.9 Although Parrini's value was measured directly, that of Kinsinger and Hughes was obtained from an exSir: trapolation to 125' of values measured at 25" and We have determined the specific refractive in- 50", and mas supported by calculations based on crement for solutions of isotactic polypropylene in the Gladstone-Dale relationship which, as the a-chloronaphthalene a t 125". Our weighted aver- authors noted, is highly unlikely to be applicable 0.005 cc./g. for to hydrocarbon polymer-a-chloronaphthalene sysage value it3 dn/dc 7 -0.189 mercury green (5460 A.) light. The measurements tems. Kinsinger'O also used a value of 1.532 for were made using a differential refractometer of the the refractive index of a-chloronaphthalene a t Debye' design as modified by Schulz.2 The in- 125". The correct value for this quantity for strument was calibrated with aqueous solutions of mercury green light is 1.594 i 0.0002, as measured potassium chloride using the data of Stamm.8 several times in our laboratory over the past The value obtained is consistent, as expected, decade, and supported by calculation from the data with measured values of dn/dc for the system poly- of Auwers and Fruhling. ethylene-a-~hloronaphthalene.~-~ It is also in Until the apparent discrepancy is resolved begood agreement with that reported by Chiang,7 tween values of dnldc reported by ourselves and -0.188 cc./g. a t 140". (Chiang's value becomes Chiang and those reported by Parrini and Kinsin-0.191 cc./g. a t 125" using a temperature coeffi- ger, it would appear advisable to apply with caucient equal to that for the polyethylene-a-chloro- tion the viscosity-molecular weight relations pubnaphthalene system, approximately $0.0002 cc./ lished for the system polypropylene-a-chlorog. "C.) naphthalene.7-9J2 However, our results do not support the value POLYCIIEMICALS DEPARTMENT N. E. WESTON of -0.216 cc./g. a t 145" reported by Parrinis (which E. I. DU PONTDE NEMOURS A N D Co., Isc. becomes - 0.220 cc./g. when similarly corrected D u PONT EXPERIMENTAL STATION
*
(1) P. P. Debye, J . Applied Phys., I T , 392 (1940). (2) G V. Schula, 0. Bodmann and H.-J. Cantow, J . Polymer Sei.. 1 0 , 7 3 (1953). (3) R F. Starnm, J . Opt. Soc. A m . , 40, 788 (1950). (4) F. W. Billmeyer, Jr., J . A m Chem. Soc., 76, 6118 (1953). ( 5 ) R. Chiang, J . Polymer Sca., 36, 91 (1959). (6) V. Kokle, .'I W. Billmeyer, Jr., L. T. Muus and E. J. Newitt. presented a t the 139th National Meeting of the American Chemical Society, St. Louis, Mo., April 23, 1961. (7) R. Chiang, J. Polymsr Sci., 28, 235 (1958). (8) P. Parrini, F. Sebastiano and G. Messina, Makromol. Chem.. 38, 27 (19601.
WILMINGTON 98, DELAWARE F. W. BILLMEYER, JR. RECEIVED DECEMBER 6, 1960 (9) J. B. Kinsinger a n d R. E. Hughes, J. Chem. P h y s . , 6 3 , 2002
(1959).
(10) J. B. Kinsinger, Ph.D. Thesis, University of Pennsylvania, 1958. (11) K. V. Auwers and A. Frilhling, Liebig'8 Ann., 422. 192 (1921), reproduced in part in the "International Critical Tables," Vol. VII, p. 49. (12) S. Shyluk, paper presented a t the 138th National Meeting of the tlinerican Chemical Society, New York, N. Y., September, 1960.
