SYSTEN ETHANOL-ISOOCTANE BETWEEN 0"
May, 1948
AND
50"
1785
From 170 to 200°, as the critical region was apTABLE 1 proached, the logarithm of the vapor pressure no SOMEVAPORPRESSURES OF PENTENE-1 FROM 0 TO 200" longer could be represented by a linear function of Temp., Vapor pressure, Temp., Vapor pressure, mm. mm. the reciprocal of the absolute temperature. The 0 241.3 100 5,221 equation, derived by aid of a central difference 5 297.1 120 7,913 table, was found to be 363.1 150 13,717 10 (4) 170-200D: loglo p (mm.) = -2089.553682 + 233,OC.
+ + 7034.8267 X
084.9289/T
low3T
-
1049.16605 X
T 2 5888.8889 X T3 Equations (1) and ( 2 ) reproduce the data to within f0.3%, while equations (3) and (4) are reliable to a mean deviation of *O.l per cent. The vapor pressures calculated by means of these equations a t a few selected rounded temperatures are given in Table I. The latent heats of vaporization were calculated for 0 and 30.07" by means of the exact Clapeyron equation ; the vapor volume was calculated by means of the van der Waals equation, the constants of which for pentene-1 were derived from critical data. These values are 6225 and (311'7calories per gram mole for 0 and 30.07'.
I C O N T R I B U T I O K FROM T H E
XORTHERX
OC.
20 531.3 760 30.07" 40 1069 60 1933 80 3269 a Normal boiling point.
19,055 22,233 25,850 27,902 30,203
170 180 190 195 200
Summary 1. The densities of liquid pentene-1 under its vapor pressure were determined from 0 to 50". 2. The vapor pressures of pentene-1 were determined from 0 to 200'; 3. Calculated heats of vaporization are reported a t 0 and 30.07". AUSTIN,TEXAS
RECEIVED JANUARY 14, 1948
REGIOXAL RESEARCH LABORATORY']
Densities and Liquid-Vapor Equilibria of the System Ethanol-Isooctane (2,2,4-Trimethylpentane)between 0 and 50 O
BY CARLB. KRETSCHMER, JANINA NOWAKOWSKA AND RICHARD WIEBE
The present investigation is part or^a systematic program in progress a t this Laboratory to determine certain physical properties of ethanol-hydrocarbon systems. A considerable amount of experimental and theoretical material on solutions of aliphatic alcohols in non-polar solvents has been published and will be referred to in its proper place. Density Measurements Commercial absolute ethanol was fractionated in a &foot column packed with glass helices and then treated with magnesium ethylate.2 The final product of dZ540.78506 was kept under its own vapor pressure in a sealed container over magnesium ethylate and samples were withdrawn by vacuum distillation. Certified isooctane (2,2,4trimethylpentanel was fractionated in the same column. The middle fraction taken was filtered through a column of silica gel to remove the small content of 01efins.~ The final density was found to be d2640.65777. The 13-ml. pycnometer shown in Fig. 1A was used for measuring the densities of the air-saturated liquids and solutions at 25". No noticeable loss of liquid through evaporation was experienced
during weighing because of the smallness of the capillaries (0.3 mm.). By applying gentle suction a t the top, filling was accomplished by means of a device shown in Fig. l B , and the turned-down tip4 greatly facilitated the adjustment of volume. The accuracy of measurement was estimated to be 2 X 10-6 g./ml. or better.
(1) One of the laboratories of the Bureau of Agricultural and Industrial Chemistry, Agricultural Research Administration, United States Department of Agriculture. (2) H. Lund and J. Bjerrum. Xu.,64, 210 (lY31). (3) A. J. Mair and A. F. Forziati, .7. R r s r w c h .Val. Bur. .Sirin~luvd%. 33, 151. 185 f J U 4 1 1
The densities of the two pure liquids at 0 and .50' relative to their values at 25' were measured in a 14-ml. Pyrex weight dilatometer (Fig. IC). Both apparatus and method were similar to those
'1;. Q-
A
Fig. 1.-A,
41
(:
F
pycnometer; B, filling device; C, weight dilatometer.
Ht.rinion
/ v i b,,.
C'hemi
A),r,iI F i
9
17'4
1
Iq i
1786
CARLB. KRETSCHMER, JANINA NOWAKOWSKA AND RICHARD WIEBE
were measdescribed by B ~ r l e w . Temperatures ~ ured by means of a standard platinum resistance thermometer and were not in error by more than 0.003°. The correction for the exposed portion of g./ml. the capillary amounted to 5 X The values obtained for ethanol are given in the last column of Table I. For isooctane the values obtained were 0.705120, 0.687773 and 0.666855 g./ml. a t 0, 25, and 50°, respectively. These values are the means of two or three determinations which differed by amounts indicating a precision of 4 X g./rnl. in the change of density over each 25' interval. In addition, the densities a t all three temperatures are affected equally by g./ml. in the pycnomthe uncertainty of 2 X eter measurements at 25' which were used to calculate the weight of liquid in the dilatometer, The densities given refer to liquids containing enough dissolved air to saturate them a t 25'. In Table I values recorded in the literature for the density of ethanol between 0' and 50' are compared with those obtained in this work. Osborne's6 values from 10' to 40' are generally believed to be very reliable. At room temperature, our value and that of Riiber' are in good agreement with them. Tyrer'ss densities are higher than the others, probably because of water in his sample. His value of thermal expansion from 0' to 25' is lower than ours, while the reverse is true for the interval 25-40', the difference in each case being of the order of mL/g. Our value of the density a t 0' is slightly higher than those of Young,YKlason and Norlin,1° and of Merriman," while our value at 50' is appreciably lower than that of Young. TABLE 1 RECORDED VALUES FOR DENSITY
OF
The following equations fit the data for ethanol and isooctane, respectively, for the range 0 to 50" Ethanol dt4 = 0.806306
Isooctane dld = 0.708120
Temp. "C. Osbornea Young6 KlasonO 0 0.80627 0.80628 10 0.79784 ,79792 20 .78934 ,7894 ,78938 26 ,78506 80 ,78078 ,78050 40 ,77203 ,7722 50 .7633
0.763137
N. S. Osborne, E. C. McKelvy and H. U'. Bearce.6 b S. Young.g P. Klason and E. Sorlin.l0 D. Tyrer.8 e R. \V. Merrirnan." I C . S . Riiber.' 0
Brooks, Howard and Crafton12have measured the density of isooctane a t 20 and 25', and our value a t 25' is in excellent agreement with theirs. S.Burlew, THISJ O U R K A L , 62, 690 (1940). (6) N. S. Osborne, E. C.hIcKelvy and H. W. Bearce, Bztl2. Bur. Standards, 9, 327 (1913). (7) C. N. Riiber, 2. Elektrochem., 29, 335 (1923). (8) D. Tyrer, J. Chem. Soc., 2534 (1914). (9) S. Young, i b i d . , 707 (1902); Sci. Pioc. Roy. Dubliiz Soc., 12, 374 (1910). (10) P. Klason and E. Xorlin, Arkizr Kemi Mineral. Geol., 2, No. 24, 1 (1906). (11) R . W. Merriman, J. Chem. Soc., 628 (1913). (12) D. B. Brooks, F. L. Howard and H. C. Crafton, J . Rescrirch A.,~L.H W . S ~ Q F Z ~ W a~4S, .3 3 (1940). ( 5 ) J.
- 8.0481
10-4t
- 3.52
X lO-*P 6.82 x
lo-Y3
X lW4t - 3.168 X 10-'t2 1.87 X 10-gt3
TABLE I1 DENSITIESAT 25 EthanolWt. fract. Mol. fract.0
1000 (V/VO 1)
-
7 -
a
ds4
0.68777 0 .68834 0.786 .68857 0.861 .69041 1.532 .69088 1.686 ,69823 2.820 ,71140 3.602 ,72735 3.806 ,74358 3.448 .76227 2.348 ,77261 1.407 1 1 ,78506 0 Molecular weights: ethanol 46.07; isooctane 114.22. 0 0.0130 ,0163 .0431 ,0499 .1432 .2960 ,4680 .63 13 .SO57 .8962
ETHANOL, 0 TO 5 0 ° ,
.77224 ,76331
- 8.4456 X
The cubic terms were chosen to make the equation for ethanol agree with Osborne's figures, and to make the equation for isooctane agree with Brooks, Howard and Crafton's value at 20°, all to g./ml. within 1 X Solutions were prepared by distilling the individual components into an evacuated bulb, and the amounts determined by weighing. In this way any loss, as well as any correction for the displacement of air, was avoided. Transfer to the pycnometer (Fig. 1A) was effected through displacement by mercury. The results shown in Table 11 were used in subsequent work for converting densities to compositions.
G./ML. Merriman' 0' Authors' Riiberf values Tyrerd 20' obtained 0.80645 0.80628 0.806306 ,79803 ,78933 ,78532 0.785063
Vol. 70
0 0.0316 .0394 .lo05 .1152 .2930 ,5104 ,6856 ,8093 .9114 .9554
Densities of solutions a t 0 and 50', referred to the density of air-saturated solutions a t 2j0, were measured with a pycnometer similar to the one TABLE I11 DENSITIESOF SOLUTIOSS AT 0 ---EthanolWt. fract.
0 0.0544 .1120 .2332 ,3232 ,4166 .5307 .5397 .6409 ,6968 ,8042 ,9195 ,9332 1
Mol. fract.
0 0.1248 ,2382 .4299 .5421 ,6390 .7371 .7440 .8157 .8507 .9106 ,9659 .9719 I
dQ4
0.70812 .71207 .71676 .72705 .73508 ,74371 ,75464 .75554 .76573 .77160 .78330 .79658 .79823 ,80631
AND
d%
0.66686 .66959 ,67378 ,68368 ,69158
,70018 ,71119 .71212 .72243 .72831 .74009 .75342 .75503 ,76314
50 O
-
1000 ( V / V o 1) 00 500
0 1.09 1.60 2.43 2.79 3.03 3.18 3.17 3.05 2.83 2.16 1.04 0.85 0
0 2.81 3.92 4.98 5.25 5.25 4.95 4.86 4.28 3.87 2.80 1.26 1.08 0
May, 1948
SYST33M ETHANOL-ISOOCTANE
shown in Fig. IA, with an estimated accuracy of 7X g./ml. Results are given in Table 111. Discussion Data of relative expansion of mixing, ( V / Vo)1, where Vis the volume of the solution and Vothe sum of the volumes of the components, given in Tables I1 and 111, are plotted in Fig. 2. Harms,13 using the mass law, calculated this volume increase on mixing for several aliphatic alcohols in cyclohexane on the basis of dissociation of complexes containing from two to an infinite number of alcohol molecules. On this basis he obtained excellent agreement between experimental and calculated values for ethanol-cyclohexane solutions, whose volumetric behavior closely resembles that of the system ethanol-isooctane. We feel, however, that the discussion is based on an oversimplified physical picture, and that the agreement is partly due to the fact that Harms' treatment contains two ad justable constants.
0
0.2 0.4 0.6 0.8 Mole fraction of ethanol. Fig. 2.--.Relative expansion of mixing.
BETWEEN
0 AND 50'
17%
the passage. Bulb J was used to remoye any possible traces of water as the ternary azeotrope. When a steady state was reached, the composition of the vapor in A was the same as that of the liquid in F, and the liquid and vapor samples were taken by means of evacuated sampling containers G. The tubes connecting the sampling containers G to the apparatus were full of air for all practical purposes and the error due to liquid holdup from this cause was negligible. No trouble was experienced using Apiezon grease L on stopcocks and joints when suitable precautions were taken.
1.0
Liquid-Vapor Equilibrium Measurements After a thorough consideration of existing methods, including actual trials, which emphasized the desirability of generating vapor in a separate boiler rather than depending on the boiling process itself to give equilibrium, the equilibrium still of Scatchard and co-workers14was selected as a basis. Purification and density df the ethanol and isooctane used have been described previously. Figure 3 shows the equilibrium still. A waterbath was substituted for the vapor jacket used by Scatchard, which simplified construction and tended to maintain A at a more uniform temperature. Vapor was generated in boiler D and passed into the vapor-lift tube B where liquid was entrained and lifted to the top of chamber A. The mixture of liquid and vapor descended through the annular space packed with 3-msn. glass helices surrounding the thermometer well, and the vapor which separated passed through tube E to a condenser and finally to trap F. A heater located at the bend prevented any condensation during (13) H. Harms, Z. pkysik. Chcm., SSB, 280 (1943). (14) G . Scatchard, el al., THISJ O U R N A L , 60, 1275, 1278 (1938); 61, 3206 (1939); 62,712 (1940); and 68, 1957, 1960 (1946).
4
Fig. 3.-Equilibrium
still: L, detail of sample bulb.
A platinum resistance thermometer and Mueller bridge, both recently calibrated, were used for temperature measurement. The condenser above trap F was connected to a 20-liter ballast and a manometer through trap K, an inverted U-tube surrounded by solid carbon dioxide. The manometer of 12 mm. i. d. tubing had provision for evacuating the vacuum arm when necessary. It was read with a Gaertner cathetometer (model M901) at a distance of 250 mm. Readings could easily be estimated to 0.02 mm., and no errors of this magnitude were found when the cathetometer was checked at the same working distance against a Gaertner standard meter calibrated by the Bureau of Standards. The vapor pressures are given in International mm. of mercury and have an estimated accuracy of 0.05 msn. After complete evacuation, dry air was admitted and approximately 52 ml. of solution was introduced into the apparatus. Distillation rates were about 35 ml./hr. at 25' and 100 ml./hr. a t
50". The psessure was regulated manually to keep the resistance thermometer reading exactly a t 25.00' or 50.00'. The bath temperature was regulated to keep a constant amount of liquid (about 3 ml.) in boiler D, which required that the bath be several hundredths of a degree above the temperature in A. During initial operation a small amount was distilled into J to eliminate any possible traces of water, as explained previously. The steady state was maintained for considerable time after which temperature and pressure were measured in quick succession, heaters turned off, the apparatus was brought to atmospheric pressure, and samples were taken. In addition, static measurements were made of the total vapor pressure of the solutions a t 0 and 23", by use of a vapor-pressure cell equipped with a magnetic stirrer and connected directly to the manometer. The solutions were freed of dissolved gases and traces of water by slowly distilling off part of the sample while the stirrer was in operation. This process was continued until McLeod gage readings on the portion of distillate uncondensed a t - 78" showed no more gases were being evolved. Vapor pressure and density of the remaining solution were then measured.
TABLE IV STATICVAPORPRESSURE MEASUREMENTS Mole fract. CzHsOH
mm. 00
p, mm. 250
0.0000 .0000 .0000 Av. .0186 .1470 .2967 .3795
13.03 13.06 13.04 19.95 22.30 22.61 22.68
49.29 49.33 49.31 78.83 92.81 95.32 95.83
0.0000 .0565 ,1182 ,1700 .2748 ,3773 ,5416 .7225 ,8511 ,9603 .9757 1.0000
00
p, mm. 250
22.65 22.18 19.94 17.99 13.81 11.96 11.94 11.95
96.05 94.41 86.31 79.64 65.28 59.01 59.04 59.02
P, mm.
104 A log P
0.0000 .4441 .4762 .4910 .5073 .5153 .5285 .5501 .5994 .7471 .8023 1.0000
49.31" 86.56 91.81 93.57 95.22 95.85 96.14 95.25 91.49 75.71 70.41 59.03"
0 - 34
in 4 1
2 6 6 27 - 70 - 120 - 16
50 '
90
80 ei
3 70 I
E 60
i (I
1.0
p, mm.
25"
r--T-------
0.2 0.4 0.6 0.8 Mole fraction of ethanol. Fig. 4.--Vapor pressure at 25".
0.5684 .7749 .go77 .9458 .9882 1.0000 1.0000 1.0000 Av.
Liquid, mole Vapor, mole fract. CnHaOH fract. CzHsOH
are presented in Tables IV and V, respectively. The good agreement between static and dynamic vapor pressure measurements a t 25' is indicated in Fig. 4; the actual agreement is within 0.2 mm.
0
Mole fract. CZHKOH
TABLE V EQUILIBRIUM-STILL MEASUREMENTS
Results Static vapor pressures a t 0 and 25" and liquidvapor equilibrium measurements a t 25 and 50'
50
p,
0.0000 0.0000 .2938 .0113 .0340 .4238 .0579 .4752 .5254 .1240 ,5701 .3428 .5176 .5863 .5943 .5941 .5969 .6144 ,7713 .6279 .6881 .8799 .7526 ,9319 .7942 .9516 .9008, .9829 1,0000 1.0000 Static measurements.
146.47 207.31 250.15 271.87 296.29 315.21 318.26 318.75 318.82 315.10 301.38 282.86 271.27 242.85 220.94
0 32 16 13 21 21 27 29 29 30 43 30 22 - 23 - 50
The vapor curve below 0.4 mole fraction ethanol is conjectural. The data for 50" give curves of similar shape The azeotropic mole fraction of ethanol a t 25' is 0.5270 and at 50' is 0.5941; the corresponding vapor pressures are 96.1 and 318.8 mm. Vapor pressures for isooctane reported here are in good agreement with the ones calculated from the equation published by Willingham and coworkers,l5as shown in the following comparison at
0, 25 and 50" with their values in parentheses: 13.04 (12.99), 49.31 (49.34) and 146.47 (146.51). Since the value quoted by Willingham, et al., for 0" represents an extrapolation from their lowest experimental point of nearly 25', the agreement demonstrates the suitability of the Antoine equation used by them. A comparison of the vapor pressures of ethanol a t 0 , 2 5 and 50" reported here with those of Merrimanll in parentheses: 11.95 (12.0), 59.02 (59.0) and 220.94 (222.2) shows a good agreement except a t 50'. Recorded values for the vapor pressure a t 50" range from 219.8 mm.16 to Merriman's
(15) C. B. Willingham, W. J. Taylor, J. M. Pignocco and F. D. Rossini, J. Reseorch N o t . Bur. Standards, 86, 219 (1945).
A177, 123 (1886).
(16) W. Ramsay and S. Young,
Phil. Trans. Roy. SOC.London,
Slay, 1948
SYSTEM ETHANOL-ISOOCTANE
222.2 mm. and the value given here is in reasonable agreement with the one of Scatchard and Raymond," via., 221.17 mm.
BETWEEN
0 AND .ioo
TABLE VI1 SMOOTHED VALUES OF THERMODYNAMIC FUNCTIONS IN CAL./MOLEAT 25 a Mole fract.
Discussion
ethanol
FE
0.05 .1 .2 .3 .4 .5 .6
93 159 250 307 335 342 325 287 225 136 76
The partial molal free energy equation for binary solutions
has been used extensively in the form of the Duhem-Margules equation
.7 .8
or substituting total pressure times the corresponding mole fraction in the vapor phase Py and P (1-y) for the partial pressures, the equation becomesla
1789
.9 .95
- TSE 28 53 100 138 166 183 183
168 138 87 49
HM 65 106 150 169 169 159 142 119
87 49 27
dicated and KM is the heat of mixing. They are shown in Table VI1 and Fig. 5 . Detailed exposi-
d l n P = 'LxdIny, 1-Y
where deviations from ideal behavior of the vapor phase are neglected. This equation was integrated numerically, using our experimental values of mole fraction of ethanol in the liquid, x, and in the vapor, y. The differences A log P between observed values of log P and those resulting from the integration are given in Table V. These deviations are made up of the experimental error plus the correction for vapor imperfections. The latter correction is proportional to the vapor pressure; a t the low pressures involved in this work it is comparable in magnitude to the contribution of experimental errors to the integral. Hence, no significant evaluation of the parameters in the equation of state of the vapor can be obtained from the listed values of A log P. The agreement between calculated and observed pressures is reasonably satisfactory, however, since only two values of Alog P exceed 0.005. Liquid-vapor compositionsa t 0 ' given in Table VI were calculated from the data in Table V on the assumption of additive specific heats. Using the same assumption, the differences F - Fi = FE and S - Si = SEas well as HMwere computed, where FE and SEare the amounts of free energy and entropy above that of the ideal solution as inTABLE VI CALCULATED LIQUID-VAPOR EQUILIBRIUM AT 0
Liquid
0.01 ,025 .05 .10
.15 .20 .30
Mole fraction ethanol Vapor Liquid
0,2673 ,, 3626 ,4062 ..4219 4291 4374 I4447
0.40 .50 .60 .70 .80
.90 .95 Azeotrope
Vapor
0.4491 ,4527 ,4592 .4698 .4952 ,5511 ,6206 .451
(17) G . Scatchard and C. L. Raymond, THIS JOURNAL, 80, 3099 (1938). (18) W. K. Lewic. and E. V. Murphree, ibid., 46, 1 (1924).
300
ai
5
200
-
i%
0"
100
/ 0
I
I
u
0.4 0.6 0.8 Mole fraction of ethanol. Fig. 5.-Thermodynamic functions at 25'. 0.2
1.0
tion of the rhethod is given by Scatchard.14 No correction was made for the imperfection of the vapor since i t was found insignificant. The system carbon tetrachloride-methano1l4 shows a practically identical shape of curve for SE os. x as well as negative values for TSEover the entire range. As pointed out, this must be due, in part a t least, to the strong interaction of the two components. The entropy of mixing was discussed recently by WOodlg a t some length and it was stated that the orientational distribution is the principal factor. The negative values of SE obtained in this investigation, as well as those obtained by Scatchard and co-workers for solutions of methanol in carbon tetrachloride and benzene, would then be explained by increased orientation of the non-polar solvent molecules caused by the presence of alcohol molecules. Such an interaction is fairly plausible for carbon tetrachloride which contains chlorine atoms that can interact with the hydroxyl hydrogen, but it is somewhat surprising for the hydrocarbons, benzene and isooctane. (19) S.
E . Wood, J . Chem. Phys., 16, 358 (1917).
the equations fail to reproduce the rapidly increasing activity coefficient of ethanol. A better fit can be obtained only by using equations with more than two adjustable constants. However, the Van Laar equations would be useful in extending data on other hydrocarbon-alcohol systems, provided the peculiar behavior a t low alcohol concentrations were kept in mind.
Summary Densities of ethanol, isooctane and of isooctane-ethanol solutions were measured a t 0, 25 and 50'. Equations giving the density of ethanol and isooctane as a function of temperature are presented. The volume expansion on mixing in0 0.2 0.4 0.6 0.8 1.0 creases rapidly with temperature. Mole fraction of ethanol. Static vapor pressures a t 0 and 25' and liquidFig. 6.-Logarithrns of activity coefficients at 50': circles, experimental values; curves, calculated from Van vapor equilibria a t 25 and 50' were determined, Good agreement was obtained between the two Laar equations. sets of measurements a t 25'. Satisfactory agreeFigure 6 shows the experimental activity coeffi- ment was also obtained when calculating vapor cients a t 50' compared with curves calculated pressures by means of the Duhem-Margules from the Van Laar equationsz0fitted to the azeo- equation without corrections for imperfection of tropic composition and pressure. The two-con- the vapor. The excess thermodynamic functions FE, TSE stant Margules equations would give nearly identical curves since the terminal activity coefficients and P were computed. Activity coefficients are nearly equal.20 The agreement is seen to be were calculated using the Van Laar equation and fairly close except below 0.1 mole fraction where approximate reproduction of the experimental data were obtained. (20) H C. CarlTon and A. P. Colburn, I n d . Enn. Chem., 34, 581 PEORIA 5. ILLINOIS
(1942).
RECEIVED NOVEMBER 13, 1947
[CONTRIBUTION FROM THE THOMPSON LABORATORY OF THE PHILLrPS EXETER ACADEMY]
Melting Point Curves of Optical Isomers B Y CHARLES L. BICKELAND ALFREDT. PEASLEE,JR.'
The problem of the melting points of mixtures of ors were working with a ternary mixture and not optical isomers was placed on a soun'd theoretical with optical opposites. basis by Roozeboom2who showed that three types The present study of the dextro and levo forms of melting point curves might be expected. Sev- of @-benzoylhydratropicacid indicates that these eral of the substances ~ t u d i e dgive ~ , ~ the mixed- optical opposites give a simple mixture. The idencrystal curve predicted by Roozeboom, a continu- tity of the isomers appears to be definitely estabous curve joining the melting points of the two op- l i ~ h e d ' ~so~ +that ~ the question involved in the tical isomers and a straight line in its simplest case of pinene should not be raised in this case.'" form. Most of the compounds i n v e ~ t i g a t e d ~The ~ ~experimental ~~ data for the acids are presented give a curve with two minima and a maximum, in- graphically in Fig. 1. dicating the formation of a racemic compound. This study has been extended to include the Ross and Somerville4 reported that pinene gave methyl esters of the above acids. Figure 2 indithe third type of curve, characteristic of a simple cates that a racemic compound is formed. The mixture and consisting of two parts with a mini- behavior of the methyl esters therefore resemmum a t the point of intersection. However, bles that of most of the acids and esters previTimmermanss stated that the "dextro" pinene ously investigated. used by Ross and Somerville was a mixture of the (7) Bickel, THIS JOURNAL, 60, 927 (1938). two isomers of a-pinene, so that these investigat(8) Kohler and Bickel, zbid., 68, 1531 (1941). (1) A senior in the Phillips Exeter Academy during the school year, 1947-1948. (2) Roozehoom, 2. physik. Chcm., 18, 494 (1899). (3) Adriani, ibid., 88, 467 (1900). (4) Ross and Somerville, J. Chcm. S o c . , 2770 (1926). ( 5 ) Ross, ibid., 718 (1936). (6) Timmermans, Bull. SOC. chim.. Belg., 39, 243 (1930).
(9) Bickel, ibid., 68, 941 (1946). (10) Since the submission of this manuscript the observations of Singh and Tewari [Proc. Indian Acod. Sci., %SA, 389 (1947)1 regarding 3-nitro-$-toluidinomethylenecamphorhave come to our attention. The d and 1 isomers of this substance appear to form a simple mixture, the melting point of the eutectic being only 1.8' below the melting point of each pure isomer.
