DIPOLEMOMENTS FOR POLAR LIQUID HYDROCARBONS
Jan. 5 , 1955
TABLE 111 TEMPERATURE DEPENDESCE OF -1im
Sucrose hydration
-1im
Maltose hydration
6-0 Cc./ Mole/ A/c g. OH 20' 0 . 1 5 0 . 2 5 0 . 6 0 0.17 25' ,1? .20 ,48 ,13 3.50 .08 ,I3 .31 ,09 450 .05 ,os ,19 .07 AH,kcal./mole 13
c-0
A/C
THE
HYDRATION Glucose hydration
- lim
Cc./ Mole/ 6-0 g. OH A/c 0 . 2 8 0.67 0 . 2 5 .22 .53 . 2 1 ,15 .36 ,17 ,12 .70 ,I3 12
Cc.1 Mole/ g. OH 0.42 0.84 .35 .70 .28 .56 .22 .44 13
From Table 111, it is found that the hydration of sucrose decreases from 0.25 to 0.08 cc.,lg. as the temperature rises from 20 t o 45". Pryor and Roscoe3 calculated the fractional volume occupied by the solute as a solvated sugar molecule, c2, from viscosity data. Such values of ce were larger than the values of the fractional volume c1derived from that of dry sugar. The value of ratio c ~ Ifor c ~ sucrose decreased from 1.23 to 1.06 as the temperature rises from 20 to 70". This change corresponds t o the fact that bound water decreases from 0.15 to 0.04 cc.,'g. as the temperatures rises from 20 to 70". In the case of glucose, the results of Roscoe gave the values of hydration decreasing from 0.10 to 0.02 (cc./g.) with the temperature rise from 25 to 50". Our results differ a little from these values. But considering the fact that different assumptions have been used in these methods, the agreement of the values may be satisfactory.
lCoNTRIBUTIOS
FROM THE
73
If we apply the theory of Langmuir's adsorption isothermg in the present case, the enthalpy change of hydration could be obtained by the equation
e
In __ = AH/RT 1 - e
+ const.
(2)
in which 8
= amount of hydration in mole number per OH
radical AH = enthalpy difference between the state of hydrated water and non-hydrated one
From those curves, we obtained the values of AH given in Table 111. These values of A H involve several other effects than the hydration: intramolecule structural changes of solute, breaking of the hydrogen bond between water molecules, etc. Therefore, the values of Table I11 cannot be discussed in more detail, but it may be mentioned that these values are of the correct order of magnitude as the energy of hydrogen bonding in all saccharides examined, and i t is evident that hydrogen bonding plays an important role in these changes. The author wishes to express his thanks to Prof. I. Sano and Assist. Prof. Y. hIiyahara for their kind advice and encouragement. (9) R. H. Fowler, "Statistical hlechanics," Cambridge University Press, 1936, p. 529.
NAGOYA, JAPAS
FRICKCHEMICAL LABORATORY, PRINCETON UNIVERSITY 1
The Determination of Atomic Polarizations and Dipole Moments for Slightly Polar Liquid Hydrocarbons',' BY
ANTHONYJ. PETR03A N D
CHARLES
P. S M Y T H
RECEIVED JLZY 10, 1957 Dielectric constants, densities and indices of refraction a t five wave lengths have been measured for the pure liquids benzene, toluene, o-xylene, m-xylene, p-xylene, ethylbenzene, styrene and isopropylbenzene a t 20, 40 and 60". The electronic polarizations have been calculated by the Lorentz-Lorenz and Cauchy relationships and the total polarizations by the Clausius-Mosotti equation. The electronic polarizations have been found to be density dependent but a plot of the difference between total and electronic polarizations against the reciprocal of absolute temperature has been found to yield a straight line whose intercept is the atomic polarization and whose slope is proportional to the dipole moment. Values obtained by this method using the Debye equation agree well with those obtained by microwave dielectric constant measurements for all except o-xylene and compounds of higher moments. The liquid and vapor oipole moments have been correlated with t h e asymmetry of the molecules.
The determination of atomic polarization (PA) for slightly polar compounds has, heretofore, been possible only by methods which are either indirect or difficult to apply experimentally or which give rather uncertain results. 4 Altshuller6 recently has determined this quantity for several liquid non(1) This research was supported by the United States Air Force through the Office of Scientific Research of the Air Research and Development Command under contract No. AF18(600) 1331. Reproduction, translation, publication, use or disposal in whole or in part by or for the United States Government is permitted. (2) This paper represents a part of the work t o be submitted by hlr. A. J. Petro t o the Graduate School of Princeton University in partial fulfillment of the requirements for the degree of Doctor of Philosophy. (3) Monsanto Fellow in Chemistry, 1956-1957. ( 4 ) A detailed summary of the various methods is given by J. W. Smith, "Electric Dipole Moments," Butterworth's Scientific Publications, London, 1955, Chapter 9. (5) A. P. Altshuller, J . Phys. Chem., 58, 392 (1954).
polar aromatic hydrocarbons but does not give values for the slightly polar compounds used in his work. It is the purpose of this paper to describe a simple method for measuring P A for compounds whose dipole moments are less than about 0.5 X Furthermore, the dipole moments may simultaneously be determined with good precision, especially for compounds whose moments are so small that they are not amenable to accurate measurement in the gaseous state or in solution.
Experimental The aromatic compounds studied were obtained from the Brothers Chemical Company, with the exception of benzene and toluene. The benzene used was the analytical reagent grade product of Merck and Company and the toluene was obtained from the Barrett Division of the Allied Chemical and Dye Corporation. The benzene and p-xylene were initially purified by fractional crystallization. o-Xylene
ANTHONYJ. PETRO AND CHARLES P. SMYTH
74
was initially purified by the method of Clarke and Taylor6 which consisted of the preparation and recrystallization of its sodium sulfonate salt followed by regeneration of the hydrocarbon by steam distillation. \Vith the exception of vinylbenzene (styrene), which was distilled under vacuum, a11 the compounds were dried over and distilled from sodium hydride or calcium hydride immediately before use. The observed boiling points arc listed in Tablc I together with literature values. TABLE I BOILIXGPOINTS OF LIQUIDHYDROCARBONS Obsd.
Benzene Toluene o-Xylene ?%-Xylene @-Xylene Ethylbenzene Styrene Isopropylbenzene
B3.p.,OC.
79.9 110.7-110.8 144.4 138.4 (747.8 mm.) 138.3-138.4 136.3 44-45 (20 nim.) 151.2
Lit.
80.1 110.6 144.4 139.1 138.4 136. 1 48 (20 rim.) 132 . 4
Vol.
so
The total or molar polarizations, P,xere calculated from and d values by the Clausius-Mosotti equation. Electronic polarizations, PE,were determined by extrapolating the molar refractions, as calculated from n and d values by the Lorentz--Lorem equation, t o infinite wave length according to the Cauchy dispersion formula. This estrapolation was accomplished by the method of least squares. Tile resulting polarization vilues are listed in Table 111. E
Discussion It has long been known that non-polar liquids show significant deviations from the ClausiusMosotti equation for the dependence of the low frequency dielectric constant on density.’O Similarly, the Lorentz-Lorenz equation has been found to give a molar refraction which varies very slightly with temperature. This fact had led t o the ernpirical equation of Eykrnan’l which, however, has not found general use. This discrepancy has been interpreted to be due t o fluctuations of configura-
TABLE 11 DENSITIES,~ DIELECTRICCOSSTAYTS’A N D REFRACTIVE INOICES~ OF LIQUIDHYDROCARBONS 1, OC.”
d
€
nmIA
~SLSIA
n m l ~
~UOIA
tl$l#A
1 50214 1.50126 1.50530 1.52308 1 49814 2 2836 1.48871 1.49273 1.51020 1 489.54 1 48597 2 2636 1 ,47586 1.49700 1 ai343 1 47688 1.47979 2 2436 1,49575 1 49303 1 49658 1.49957 1.51666 Toluene 2 3837 1,48496 1 48379 1.48888 1.50558 1 48220 2 3390 1.49438 1,47394 1 47501 1 ,47779 1 47157 2 2941 60 1.50540 1.52613 1 50014 1,50911 1 50242 2 5872 o-Xylene 20 1 19611 1 ,49930 1 ,49553 1.51606 1 49289 2 5315 40 1 ,48343 1 46641 1 50506 1 4!5297 1.48926 2 4781 60 1,51769 1.49717 1 4‘)7‘71 1 49419 1.50094 2 3684 30 1IZ-Sylcnc 1,48723 1 48YOi) 1,5073 L7 1,49091 1 48452 2 3268 40 1,49690 1,4769s 1 47795 1.45073 1 47416 2 2860 60 1.49910 1 ,48537 119611 1.a581 1 49253 2 2099 p-Xylene 20 1 48622 1.50554 1,485213 1.48912 1 48259 41) 2 23C8 I ,47483 1 47.iSCl 1.47877 1.49496 1 47265 2 2034 60 1.49575 1.49938 1.51553 1 49659 1 49286 2(1 Ethylbenzene 2 4042 1 48617 1.50530 1.48563 1.45932 1 48298 2 3587 40 1 47604 1,49448 1,47507 1.47867 1 47284 2 3136 60 1.57917 1 54503 1.55265 1 54296 1.54678 Styrene 2 4257 20 1,53561 1 53704 1.54141 1.56771 2 3884 40 1.55603 1.52459 1 52598 1.53033 2 3510 60 1 4921)2 1.49138 1.50996 1.49182 1 48858 Tsopropylbenzene 2 3833 20 1 45245 1.50016 1.48502 1.48166 1 47930 2 3386 40 1 472’36 1.47499 1.48981 1.47149 1 46930 2 2958 GO a 4Z0.0000,5. * 4~0.0002. &0.00005. 0 &0.05”. f American Petro1e:m Institute Research Project 44, “Selected J. Timmermans, Physico-Chemical Constants of Pure Organic Values of Hydrocarbons and Related Compounds.” Compounds,” Elsevier Publishing Co., Inc. * R. H . Boundy and R . F. Boyer (editors), “Styrene: Its Polymers, Copolytilers and Derivatives,” Reinhold Publ. Corp., New York, S.Y., 1952, p. 35. hfeasured values: d 2 j 0.9316, d40 0.8885. Benzene
20 40 60 210 40
0 . 8790gf .85769’ ,83602’ ,8669‘ .8483Q .83000 ,8800 1 ,86338 ,84643 .86470 ,84735 .82986 .8611“ .843V .52609 ,86099 .84949 ,83171 .9063h .8 8 G h . 8721* ,86192 ,84501 .82779
The dielectric constants, E , of the pure liquids were measured a t 20, 40 and 60” with an apparatus previously described or referred to.’ Refractive indices were measured with a Pulfrich refractometer using the sodium-D line and fnur lines in the mercury vapor spectrum. For measurements a t 40 and 60” the index of refraction of the prism at each v a v e length was determined with distilled water .8 these temperatures some scattering of the light was observed but by careful manipulation the critical ray could be sharplv defined. \%%en good literature values mere not available, densities were determined with a graduated px cnometer,g with an accuracy of =t0.002();. The experimental values are listed in Table 11. (6) H. T. Clarke and E. R. Taylor, Txixs JOURNAL, 46, 831 (1923). (7) A. J. Petro, C. P. Smyth and L. G. S. Brooker, ibid., 78, 3040 (1956). (8) “International Critical Tables,” Vol. V I I , McGraw-Hill Book Co., h’ew York. N. Y..p. 13. (9) G. R. Robertson, I n d . Eng. Chem., Anal. Ed., 11, 481 (1939).
tion and anisotropy of the molecules caused by their translational, rotational and internal motions, which give rise to microscopic changes in molecular polarizability.12 Consequently, equations have been derived to take this effect into account. The recent work of Brown13 contains a detailed comparison and evaluation of the various modifications of the Clausius-Mosotti equation. As a result, i t is apparent that either PEor both P E and P A must be variable with density. The data listed in Table I11 show PI.:to be temperature or density dependent. (10) S. S . Kurtz, Jr., and A . I,. Ward, J . Franklin Inst., 222, 563 (1938); 224, 583, 697 (1937). (11) J. F. Eykman, Rec. Iran. chim., 14, 185 (1895). (12) F. G. Keyes and J. G Rirkwood, Phys. Rev., 37,202 (1931). (13) W. F. Brown, Jr.. J . Chem. Phys., 18, 1193, 1200 (1950).
DIPOLEMOMENTS FOR POLAR LIQCIDHYDROCARBONS
Jan. 5, 1958
TABLE I11 MOLARAND ELECTRONIC POLARIZATIONS (Cc.) MATIC HYDROCARBOXS I,
Benzene
Toluene
o-Xylene
m-Xylene
p-Xylene
Ethylbenzene
Styrene
Isopropylbenzene
=c 20 40 60 20 40 60 20 40 60 20 40 60 20 40 60 20 40 60 20 40 60 20 40 60
P
PE
(10.01)
(10.01)
26.62 26.69 26.75 33.55 33.51 33.45 41.74 41.56 41.40 38.46 38.42 38.39 36.66 36.74 36.80 39.04 38.96 38.87 37.02 37.07 37.12 44.00 43.89 43.80
25.10 25.18 25.24 31.54 31.61 31.68 34.39 34.48 34.56 34.53 34.62 34.69 34.56 34.65 34.72 34.38 34.46 34.55 34.49 34.56 34.63 38.93 39.03 39.11
OF
P
ARO-
- PE
1.52 1.51 1.51 2.01 1.90 1.77 7.35 7.08 6.84 3.93 3.80 3.70 2.10 2.09 2.08 4.66 4.50 4.32 2.53 2.51 2.49 5.07 4.86 4.69
Furthermore, since, for benzene and @xylene, P varies by the same amount and in the same direction as PE,i t is evident that PA is constant within the experimental error over the temperature range studied in agreement with the behavior generally observed. l 4 The fact that the molar refraction is a function of density rather than of temperature can be seen by examining the data of Gibson and Kincaid.15 Their values of the Lorentz-Lorenz specific refraction decrease with increasing pressure and consequently decreasing specific volume a t constant temperature, but, a t approximately equal specific volumes a t different temperatures and pressures, the specific refractions are practically equal, as illustrated belowl5 f,
=c. 25 35 45
P (bars) 26 154 280
Y
(cc.)
1.14186 1.14171 1.14158
Spec. refr.
0.3478 ,3478 .3478
In the present work, the change in density is brought about by changing the temperature, but the latter variation has no direct effect on the refraction, a t least in the cases of benzene,15carbon dioxide and carbon disulfide.13 The attempt was made to utilize this effect to determine the dipole moment and atomic polarization for a pure liquid. If the Debye equation is written
- -?E = P A f 47rN~2 (1) then a plot of (P- PE)vs. 1/T might yield p from p
the slope and P A from the intercept.
The values
(14) C. P. Smyth, "Dielectric Behavior and Structure, "McGrawHill Book Co., New York, N. Y., 1955, p. 417-422. (15) R. E. Gibson and J. F. Kincaid, THIS J O U R N A L , 60, 511 (1038).
75
determined from equation 1 are listed in Table I V under PA^ and p1. I n order to minimize errors all calculations were made by the method of least squares. Also included in Table I V under PA^ TABLE IV DIPOLEhfOMENTS' ( x 10") AND ATOMIC POLARIZATIONSD (CC.) O F LIQUIDHYDROCARBONS
.. 0" 1.50 .. 0 Benzene 0 . 3 1 0.31 0.37" 1.77 1 . 7 5 Toluene .62d 3.11 1 . 9 8 .45 .52 o-Xylene .30 .31 ( .37)' 2 . 0 1 1.96 m-Xylene .02 Od 1.95 .. p-Xylene .37 .58' 1.84 1.84 .37 Ethylbenzene .13
