HAZIME SHIIO
70
indicates that the activation energy for the H atom displacement process is about equal to or slightly less than that of the H atom abstraction from crotonaldehyde. The value for the latter reaction is not known, but the similar reaction with acetaldehyde is reported to have an activation energy (ikcal." If, of course, chain A supplies a significant ainouiit of propylene the kinetics become more coniplex and this treatment breaks down.
CDS $- C H F C H C H ~ C Hd ~ [CD3CHeCHCH?CH3] [ C D ~ C H Z C H C H ~ C+ H ~ ] CD3CH2CH=CH? f CHa
A similar intermediate has been suggested in the methyl displacement of acetyl from methyl propenyl ketone3 and may exist in the analogous process with crotonaldehyde. However, there is littlc if any direct evidence for these intermediates. An interesting process, recently suggested by Kistiakowsky and AIahan'6 to explaiii their results from the photolysis of methyl ketene, is the attack of the ethylidene radical on methyl ketene to give butene and CO
summary It now seems well established that displacement processes are important modes of reaction of methyl radicals with biacetyl, l ? trans-methyl propenyl ketone and crotonaldehyde. Recently other displacement processes have been reported to occur when methyl radicals react with propylene and 1butene. VarnerinI3 states that a t least 5% of the methyl radicals generated in the pyrrolysis of mixtures of CD3CD0 and propylene react by (22)
+ CH-CHCH, ---+ CDaCH=CH2 + CHI
CHBCH f CH,CH=C=O
+ CH?=CHCHzCH3
I1
(22)
CH-CHCD,
€
+ CHFCHCH~CD~
CH-CHCH,
+ CD3
+ CHtCD3
the investigators propose that methyl radicals first add to the double bond to form an intermediate radical which subsequently dissociates into a new olefin and another free radical. For example, Mc(11) W. R. Trost, B. de B. Darwent and E. W. R. Steacie, J. Chem., Phys., 16, 353 (1948). (12) F. E. Blacet and E'. W.Bell, Disc.Faraday Soc., 14, 70 (1653). (13) R . E. Varnerin, Tins J O U R N A L , 77, 1426 (1955). (14) J. R. McKesby and A. S. Gordon, Abstracts Div. Phys. and Inorg. Chem. Spring Ileeting, Am. Chem. Soc., Miami, 1957. (15) P. Kebarle a n d I\'. A. Bryce, C a w J . Ckem., 36, 576 (1957).
[COSTRIBUTION FROM
+ CHaCH=CH2 1%
+ CrH5
CII3CHpCH=CHz
+ CO
-+-
+CI+.=CH?
CHB
+ CH3CS
HCX
+ Cllg
The recent accumulation of this body of evidence for free radical displacement processes indicates that secondary reactions of this type may be of considerable importance in a number of other systems similar to those described. Current studies in this Laboratory include an investigation of the reactions of methyl radicals with the a,p-unsaturated ester, methyl crotonate. Acknowledgments.-The authors are indebted to the National Science Foundation for a grant in support of this research. JVe are also indebted to Dr. R. S.Tolberg for assistance in assembling much of the apparatus employed in this problem. Mass spectrometric analyses were conducted a t UCLA through the courtesy of Professor F. E. Blacet, Dr. R. Holroyd and N r . R. Vanselow.
n'ithin a few weeks the reverse reactions were reported by Kebarle and Bryce.'5 In both instances CIIS
CiHP
and LVinkler17and reported by Forst and n'iiik1c.r. **
e CD3CH2CH=CH2 + CH3
-L
--+
This process can be classified as a displacement of carbon monoxide by the ethylidene radical. Two examples of what might be termed H atom displacements are proposed by Iiabinovitch, Davis
LIcXesby and Gordon photolyzed acetone-& in the presence of I - b ~ t e n e a' ~t 375 and 500" and explain their results by a mechanism which includes the displacement process CD3
so
Nesby and Gordon visualize the two step process
.
RIVERSIDE, CALIF.
THE CHEMICAL INSTITUTE,
FACULTY OF SCIESCE, KAGOYA
UNIVERSITY
1
Ultrasonic Interferometer Measurements of the Amount of Bound Water. Saccharides BY
HAZIME
SHIIO
RECEIVED JUNE 19, 1957 Ultrasonic velocities in aqueous solutions of a number of saccharides have been measured with an ultrasonic interferometer, and the amounts of hydration have been determined a t 25'. The following samples were studied: xylose, arabinose, fructose, glucose, a-methyl glucoside, sucrose, maltose, cellobiose and rafinose. The amount of bound water obtained is 0.2% 0.42, 0.38, 0.35, 0.23, 0.20, 0.22, 0.25 and 0.22 cc./g. solute, respectively. From these values, it is shown that 0.5- 0.9 water molecules are hydrated to one OH radical of each saccharide. Further investigations were carried out on the hydrations of a few saccharides a t various temperatures, and the enthalpy differencebetween hydrate water and non-hydrated 113s been estimated to be about 12-13 kcal./mole.
Introduction The compressibility of the solution may be determined by the effects from solvent, solute and solvation. The effects of the solute are separated
into two parts: the compressibility of the solute molecule and solute-solute interaction. If the concentration of the solution becomes sufficiently low, the second effect becomes negligible. The
ULTRASONIC VELOCITIESOF AQUEOUSSOLUTIONS OF SACCHARIDES
Jan. 5 , 195s
compressibility of low molecular weight solute molecules may be negligibly small compared with other effects. I n the previous report,’ we have shown by adding the ethanol to the solution that the compressibility of the sugar molecule was nearly zero. I n such case, therefore, the effects of the solute should not be taken into account, and we can obtain the amount of hydration from measurements of the compressibility of solvent and solution. Passynsky2 estimated the hydration of sucrose by measuring the ultrasonic velocity, but his method was a little different from ours. Pryor and Roscoe3 measured the ultrasonic velocity of the solutions of sucrose and a few other saccharides a t temperatures from 20 to 80” and showed qualitatively that “solvation envelope” decreased with increasing temperature, provided the sugar molecules could be regarded as incompressible. But they did not extrapolate the measurement to zero concentration as we did, and the effects due to the interaction between solute molecules still existed a t the concentrations of their experiments. In this report, employing previous method, the hydrations are investigated for a number of saccharides over a wide temperature range. Experimental The ultrasonic velocities in solutions were measured with an interferometer using X-cut crystal of resonance frequency 1 mc. The materials were E.P. grade chemicals, and were dried in a vacuum desiccator for a week. The measurements were made a t 25’ for solutions of most of the saccharides, and for glucose, sucrose and maltose a t 20, 2 5 , 35 and 45”. The methods were described in a previous r e p ~ r t . ~
0.00 -
6. -0.05
-
d
0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 Concentration (g./cc.) Fig. 1.-Sucrose solutions a t 20”.
is proportional t o the volume of bound water. The value of /32 in the proportionality constant (1 - p2/po) has been arbitrarily assigned by various authors; some gave it zero value for both electrolytes2J and n o n - e l e ~ t r o l y t e s . ~It ~ ~may be justified in electrolytes owing to strong electric field surrounding ion. However for non-electrolytes, such as saccharides, the change of volume in dissolution process is negligibly small, so that the difference of the compressibility between hydrated water and normal water may be smaller than in the case of electrolytes. There are several methods to obtain the bound water406: that is, vapor pressure
Results and Discussion According to the previous report,’ neglecting compressibility of the solute, we have lim 4 l c C-bO
= @/PO -
VO)/C = -tdl
- PdPd/c
71
0
(1)
The notations used are P
= adiabatic compressibility of s o h . PO = adiabatic compressibility of solvent PZ = adiabatic compressibility of bound water YO = specific vol. fraction of bound water VO= apparent specific vol. fraction of total solvent used * . e . , VO= ( d ,- c ) / d o d = density of solution do = density of solvent c = concn. of solute, g./cc. of s o h .
0 -Sucrose
By the measurements of p, PO, Vo and c, we obtain A/c a t various concentrations of saccharides. The results and their values extrapolated to zero concentration, lim A/c, are shown in Fig. 1 and c-0
Table I. The experimental values of A/c for sucrose, Fig. 1, increase linearly with concentration, probably by reason of the solute-solute interaction depending on the square of concentration. Therefore, the values of lim A/c in Table I were obtained graphiC+O
cally by the extrapolation. AS is seen in equation 1, the value of -1im
A/c
C+O
(1) H. (1955). (2) A. (3) A. (4) H.
Shilo, T.Ogawa and H. Yoshihashi, THISJOURNAL, 77, 4980
0.0033 0.0034 l/T. dependence of the hydration.
0.0032 Fig. 2.-Temperature
lowering or freezing point depression methods determining “non-solvent water,” calorimetric or dilatometric methods of “non-freezing water,” cobaltous chloride method, dielectric method, etc. I n these methods, it is mentioned that the bound water is of solid-like structure to lose the freedom of mobility, and the free volume contained in the free water is decreased ; then, the compressibility of bound water may be considerably smaller than in ( 5 ) T. Yasunaga and T. Sasaki, J. Chrm. Sac. Japan, Pare Chem.
Passynsky, Acto Phyricochim. U.R.S.S., 22, 137 (1947). Pryor and R. Roscoe, Pmc. Phys. Soc., B67, 70 (1954). Shiio, KQgQklL%O RyOik:, J a p a n , 11, 440 (1957).
S a . , 72, 366 (1951). (6) F. Haurowitz, ”Chem. & Biol. of Proteins,” Academic Press New York, N. Y., 1950.
1-01, so
72 TABLE I EXPER:IMESTAL C
(g icc.)
Xylose
Araliiiirise
Ihictosc
(;liiro.;e
a-Methyl glucoside
Sucrose
X I alt m e
Celiol,iose
IZafittrise
RESUI '.TS AT 26' o/CJ.
d 1.0679
- l i m I.',
1'0
-- 3 / c
0.1994 0.8505 0.8791 0.126 ,1669 1.0574 ,8728 ,8931 ,122 ,9064 ,1193 1,0404 ,9238 ,146 1010 1.0339 ,9350 ,145 ,9903 ,0918 1.0300 ,5416 ,158 ,9271 0768 1.0249 9397 ,14509 , 14h .I979 1.0714 ,8382 .87C1 .l!32 1573 1 ,0560 .8fjS1 ,9013 211 I 260 1 .044(i ,8R4A .RL'IR ,213 1052 1.03Gti .9.341 ,9095 ,234 ,0853 1.0303 ,9248 .9448 ,226 OF72 1 . 0 2 7 1 ,9431 ,9577 ,217 ,1835 1.0076 ,8328 .88GG ,184 ,1390 1.0306 ,8800 ,9143 ,204 ,1349 1 ,0490 ,8883 .91G8 ,210 ,1199 1 . 0 4 3 3 ,9019 ,92r;1 ,202 ,0865 1.0305 ,9277 ,9458 ,209 ,1971 1.0716 ,8455 ,8770 .1GO . 172.6 1 . 0 6 2 2 ,8644 , 8 9 2 2 , 1 6 1 ,12?6 1.0436 .go27 ,9237 ,181 .a597 1.0196 ,9513 ,9628 ,194 .I421 1.0422 .8676 .90Yi .I06 .os94 1.0286 ,9191 ,9319 ,128 .9.%61 ,123 ,9364 ,0787 1.0220 .0724 1.0201 ,9422 ,9304 .113 ,0640 1.0174 ,9180 ,9562 ,128 ,1585 1.0377 .0018 .OS3 ,8683 1 4 2 6 1.0518 ,8987 .0118 ,092 ,1234 1.0443 ,9127 ,9237 ,089 ,0979 1.0346 ,9302 ,9395 ,096 ,9379 ,9463 ,100 .O8ijG 1.0303 ,1760 1.0646 ,8738 ,8912 ,099 .I276 1.0463 ,9072 ,9214 ,111 1199 1.0433 .921,1 ,9129 ,110 .0730 1.0252 ,9463 ,9550 ,119 ,0421 1.0133 ,9751 9688 .I25 , 1 6 5 2 1.0603 ,8624 ,8977 ,115 ,1260 1.0-133 .on54 ,9220 ,140 ,1002 1.0354 ,9252 ,9379 ,127 ,9471 144 ,08:15 1.0298 ,9348 ,0765 1 . 0 2 6 2 .9425 ,9525 ,131 .0677 1.0229 ,9492 .9580 ,130 ,9644 ,9711 ,143 ,0407 1.0150 1314 1.0486 ,9050 ,9199 ,113 ,1149 1.0417 ,9169 ,9296 ,111 ,1011 1.0364 ,9380 ,110 ,9209 ,9471 ,121 .0862 1 , 0 3 0 5 ,9367 ,0729 1 0254 ,9466 ,9553 ,119
c-0
0.17
', -, _.
,2R
,',I
$14
.I2
,1R
.l5
,13
normal state (the compressibility of water and ice are about 45 and 18 X 10-l2 cm.Vdyne, respectively). Thus we have used 18/45 for the value of &/Po, and a reasonable value has been obtained for the hydration of dextrin in a previous report.' Using this value for &/Po, the hydrations of saccharides are found as in Table 'IT. TABLE I1 ROITNn n r A r F R OF SACCHARIDES Cc ' g
Xylose .lrabinoie Fructose Glucose ol-Meth~lgl~tcositl(.
Sucrow Maltose Cellobiose Raffinose
0 28
42 3s 35 23 20 22 25 22
(25')
\Iole/mole
2 3 3 3 2 3 4
3 5 8 5 6
8
2 4 8 6 2
Mole/OH radical
0.58 89 76 .70 63 48 .53 60 -56
The hydrations of sugars in solution have been investigated by a few authors who used various
methods. Roscoe3 calculated the ratio of the fractional volunie o f solvated solute to thc dry sol ute from viscosity data, finding that the value of the h y c l r ~ t i ~of~ isucrose was 2.7 inolc mole solute a t 20". Miller' obt:iined the hydration of nun-electrolytes from d,tta of the inolality activity coefficients, using the method of Stokes atid Ikr\
.iii(l
I\
\
l < c ~ I ~ , n w i 7i
ill',
101 K V A I . , 70. I8711
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 the 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
