T H E LATENT HEAT O F FUSION OF KAPHTHALENE FROM YEW SOLUBILITY DATA* BY ARTHUR A. SUNIER AND CHARLES ROSENBLUM
The method of Hildebrand and Jenks' for evaluating solubility data has been applied by Mortimer* to binary solutions involving non-polar and in several instances even polar components. The method consists in plotting the logarithm of the mol fraction of solute against the reciprocal of the absolute solubility temperature. From the ideal solubility equation LogN =
-
L
-+ I 4.58 T
where L is the molal latent heat of fusion of solute, N the mol fraction of solute, T the absolute solubility temperature, and I a constant of integration, it follows that for an ideal solution, the curve obtained is a straight line Log N the slope of which is given by S = _ _ . For most solutions involving 1
IT
non-polar components, a fairly representative straight line can be drawn through the points corresponding to higher concentrations of solute, and a slope obtained. If substances of distinctly non-polar nature are used, assuming that no molecular changes take place in solution, the slopes are indicative of the relative internal pressures of solvents concerned. hlortimer has constructed a nomograph with a great many compounds arranged according to their relative internal pressures and finds that by employing factors representing the ratio between experimental slopes and the ideal, he can predict the solubility of these substances in one another with a fair degree of accuracy. As the basis for relative internal pressure comparison Mortimer uses that of naphthalene, which gives an ideal slope of 970 as calculated from the value 34.69 calories per gram given by Bogojawlenski3 for the heat of fusion a t its melting point. The more recent work of Andrews, Lynn, and Johnston' on the latent heat of fusion of organic compounds indicates a somewhat higher ideal slope of 990 for naphthalene. However, Ward,j who has made an extended investigation of the solubility relations of naphthalene, has obtained a slope of 970 from the upper portion of the Log lu' versus I / T curve for the system naphthalene-chlorobenzene which appears to be approximately ideal. * Contribution trom the Chemical Laboratory of the Cniversity of Rochester. 1
J. Am. Chem. Sac., 4 2 , 2180 (1920). J. Am. Chem. Sac., 4 4 , 1416 (1922); 45, 633 (1923). Chem. Zentr., (5) 9 11, 945 (1905). J. Am. Chem. Sac., 48, 1274 (1926). J. Phys. Chem., 30, 1316 (1926).
I050
ARTHUR A. SUNIER AND CHARLES ROSENBLUM.
It seems likely that a study of solutions which do not deviate greatly from the ideal would throw light upon the magnitude of the heat of fusion of naphthalene. Mortimer' has calculated, apparently from freezing point data, the slopes of the Log N versus I/T curves for naphthalene in ethylene chloride and bromide, and finds those solutions almost ideal. Accordingly the solubility of naphthalene in these solvents as well as in the corresponding ethylidene compounds has been determined in an effort to obtain a more definite indication of the latent heat of fusion of naphthalene from solubility data.
Experimental The synthetic method of Alexejew2 has been employed in making the solubility determinations recorded. The method is essentially the heating of weighed quantities of solute and solvent in a sealed tube and noting the highest temperature a t which but few small crystals of solute remain in equilibrium with solution. Precautions assuring the attainment of equilibrium a t the solubility temperature have been described by Ward. A bath of four liter capacity, a heating rate of not more than a hundredth of a degree per minute, and rotation of the sealed tube by motor were found to give satisfactory results. Uniformly small crystals were first obtained by heating to above the solubility point and rapidly cooling. The maximum error in the solubility temperature is estimated by Ward to be not more than three tenths of a degree. Ward's procedure, except for certain modifications, has been used in this investigation. The water bath was heated by means of Cenco lagless heaters or Hotpoint immersion heaters in series with resistance sufficient to cut out all current if so desired. The resistances allowed accurate control of the rate of temperature increase. The procedure in heating was somewhat modified. Instead of heating at a uniformly slow rate, an attempt was made to approximate conditions of the so-called equilibrium method of determining solubility. The temperature was raised rapidly to within several degrees of the solubility temperature, then more slowly a t about a tenth of a degree per minute to within three or four tenths of a degree from the solubility point which can be approximated from the general appearance and quantity of crystals in the tube. The bath temperature was then kept constant at each tenth of a degree for about an hour till the solubility point was reached. The recorded solubility temperature was the highest temperature a t which a very small quantity (several small crystals) of solute remained in equilibrium with solution. Immediately upon recording the solubility temperature, the bath was cooled a t a rate of about 0.05 degree per minute. I n all determinations recorded, marked increase in size of the crystals was noted within 0.1' to 0 . 2 ~ . This is important as an indication that super saturation did not take place and that the solubility temperature as noted represents equilibrium conditions. Beckmann 1
J. Am. Chem. SOC., 44, 1416 (P922).
* W e d . Ann., 28, 30j (1886).
LATENTHEATOFFUSIONOFNAPHTHALENE
1051
thermometers were used to check the rate of increase and the constancy of the bath temperature. This modification in heating procedure was made in an effort to overcome the error introduced by differences in the size of crystals. Pyrex glass tubing of 7-9 mm. inside diameter and a half millimeter or less thick was used for the solubility tubes. The heating procedure tends to overcome the difficulty raised by the time required for the solution to attain thermal equilibrium with the bath through the glass wall. I n making up the solubility tubes an attempt was made to restrict the vapor space to a very small volume. For tubes containing low naphthalene concentrations, it was found advisable to add the solvent first. The solid component formed a sort of protective layer preventing excessive evaporation of solvent and allowing the tube t o be sealed closer to the surface of the material. For higher naphthalene concentrations the solute was added first, the greater part of it being fused in the tube by immersion in a water bath between 85 and 90' C. The fusion was carried out with the solubility tube inside a stoppered eight inch test tube to prevent absorption of water vapor. Uniform bath temperature was insured by a twelve inch Cenco turbine stirrer (1000R.P.M.). The contents of the solubility tube were shaken by attaching the tube to a cross-piece one end of which was alternately raised and lowered by an eccentric motion. It was found the bubble of vapor moving from one end of the tube to the other agitated the crystals of solute sufficiently. The bath was surrounded by a wooden casing lined with asbestos pads to prevent any cooling effect from drafts. The accuracy of the method was checked by repeating some of Ward's n-ork on benzene. The benzene was purified partially according to the method of Richards and Shipley' except that no recrystallizations of solid benzene were attempted and distillation was carried out over phosphorus pentoxide. It was found that Ward's data could be checked within the error estimated for the method. Some of the work on naphthalene-ethylene chloride was repeated using somewhat poorer materials. The deviation was well within the three-tenths of a degree limit. Some preliminary experiments on more rapid heating rates were carried out. It was found that heating as rapidly as a tenth of a degree per minute when approaching the solubility temperature occasioned only seldom an error of three-tenths of a degree. Measurements were not carried below room temperature since, as Ward has pointed out, deviations from even an approximate straight line relationship are always greater a t lower temperature. The thermometers used were compared with a thermometer recently standardized by the Bureau of Standards. It is believed that temperature measurements were never in error as much as +0.07 J. Am. Chem. Soc., 36, 182j (1914).
I052
ARTHUR A. SUSIER AND CHARLES ROSENBLUM
Materials All materials used were obtained from the Eastman Kodak Company. Two samples of naphthalene were used in the determinations. Both were purified by recrystallization from methyl alcohol dried over calcium oxide and distilled. A small quantity of sample A was recrystallized thirteen times making small cuts, whereas a large amount of sample B was recrystallized eight times making very large cuts in the amount of solid crystallizing out. Both samples were used interchangeably in some of the determinations without any apparent deviation from the temperature versus concentration curves. The purity of the naphthalene was tested by taking its melting point according to the method of the Bureau of Standards.' The final fraction of sample A held over one tenth of a degree for more than forty minutes using an ordinary thermometer graduated to degrees. The fraction of sample B tested (after five recrystallizations) gave a melting point of 80.10' to 80.12' using the prescribed naphthalene thermometer obtained from the Taylor Instrument Company. Yo attempt was made to purify excessively any of the solvents used other than by fractional distillation through a Clarke-Rahrs column. Cohen2 has shown that the solubility of anthracene is little influenced by traces of moisture. The same probably holds for naphthalene, a very similar compound. Each solvent was twice distilled and representative fractions from the second distillations were used in the final determinations. Preliminary work with ethylene chloride only once distilled showed no deviations from the final results herein recorded. The corrected boiling points of the solvents used are:Ethylene chloride . . . . . . . . 8 3 . 4 3 to 8 3 . 4 8 C. Ethylene bromide3 , , . . . . . . 131.1 (131.5)(131.7) Ethylidene chloride Ethylidene bromide
. . , .
. . . .
. . . . . . . .
5 7 . 3 to 5 7 . 4 106.8to 107 . o
Results The experimental results have been incorporated in Tables I-IV. Table T' contains the solubility of naphthalene in all four solvents a t rounded temperature. All experimentally determined points have been plotted in Figs. I and 2 . As is apparent from the figures, the points are too close together to allow individual curves to be drawn for each solvent on so small a scale. Sci. Paper Bur. Standards, No. 340. 2
2. physik. Chem., 119, 247 (1926).
At the conclusion of the solubility work on ethylene bromide i t was found that Timmerman and Martin (J. Chim. phys., 23, 747 (1926)) found that its boiling point was I J I . ~ " when , determined with great care. After some time it was found that the Anschute thermometer (used only for the boiling point of ethylene bromide and ethylidene bromide) was unreliable to the extent of 0.4 to 0.6". Two other samples of ethylene bromide were prepared; their boiling points, dtermined with great care, are given in the parentheses. Five new solubility tubes were made up with these two new fractions, and the solubilitx temperatures determined in the usual manner; four tubes were on the average only 0.1 low, which is very satisfactory agreement. It seems certain, therefore, that all three fractions of ethylene bromide were quite pure, but the recorded boiling point of the first is in error.
LATEKT HEAT O F FUSIOS OF SAPHTHALESE
Tanm 1 Solubility of Saphthalene in Ethylene Chloride Grams C2HaCL
Mol Fraction Cii"
o 918 o 827 o 678
2.13j I 966
o 164 0 340 0 931 I 236 1.433
1.627
1.748
Grams
CirHs
2.386 2.093 2 .527
I'
I93
0 j72
Solubility Temperature j j
ioc
69 7 59 4 51 1
1.821 TABLE 11 Solubility of Saphthalene in Ethylene Bromide
Grams
CiaHj 2.213
2.804 2 .OjI I
'
899
1.32j I . 103 I ,065
1101Fraction
Grams C?H4Br2
Cir&
Solubility Temperature
0.184 0.966 1.259 I . 740 2.062
0.946 0.810 0. joj 0.616 0.48j
44.4
2 . j I j
0.392 0.268
35,3
4.287
76.8OC 68.9 61.7 55.'
20.7
TABLE I11 Solubility of Kaphthalene in Ethylidene Chloride Grams CioH8
Grams CHEHCII
Mol Fraction CioHs
2.637
0.I80
0
2.402
0
0.824
2.081
396 0.612 0.737 I .291
I .3 2 0
I . 152
0.470
1.259
1.629
0.374
1.938 I .800
919
0.j I 0
0.654 0.555
Solubility Temperature
ij 6°C 69 8 62.6 58.6 50.8 43.3 33.0
TABLE I11 Solubility of Kaphthalene in Ethylidene Bromide Grams CiaHr
Grams CH3CHBr2
Mol Fraction CioHs
Solubility Temperature
2.231
0.370 0.540
0.898 0.807 0.726 0.606 0,545 0.496 0.31j
74.3OC 68.5 63 .o 53 ' 7 48.4 44.1 23.9
1'539 1.948 2.063 2.063 I . 189 I ,016
1.082
I ,972
2'535 1,772
3'239
ARTHUR A. SUNIER AND CHARLES ROSENBLUM
I054
TABLE V Solubility of Naphthalene a t Rounded Temperatures expressed as 1\10] Fraction Naphthalene Temperature
CHaCHBr2 0.910 0.830 0.75.5 0.685 0.621 0.562 0.506 0.456
C2HaClz 0.911 0.831 0.756 0.686
75OC 70 65 60 55
0.620
0.560 0,503 0.452
50
45 40 35 30 25 20
0.405
0.409
0.360 ( 0 . 3 18)
0.366 0.324 (0,286)
C2H& 0.910 0.828
CHICHCI? 0.909
0.752
0.745 0.674 0.608
0.682 0.61j 0.532 0.493 0.439 0.389 0,344 0.302 (0.262)
0.825
0.545 0.488 0.437 0.391 (0.3501
28 3.0
3.2
7.4 4
7.L
i.6
7.7
T8
T.9
LOG MOL FmAcrioN NAPTNALZNE
FIG.I
Discussion
As has been pointed out, the slope of the straight line obtained by plotting the Log N versus r/T curve for an ideal solution is given by S = L/4.58T. For solutions of non-polar components deviations from this slope are indicative of the relative internal pressure of solvent. For solutions which approach the ideal, the upper portion of the Log X versus I / T curve should be a straight line with a slope related to the latent heat of fusion of solute. For the solutions herein investigated, Table VI contains the slopes indicating heat of fusion of naphthalene and the range through which the straight line function holds. The table seems to indicate that the lower value for the latent heat of fusion of naphthalene is the more likely. Except for the case of ethylene dichloride, which gives a slope somewhat less than 970, the solutions in ques-
LAI'EN'I' HEAT OF FUSIOS OF NAPHTHALESE
'055
tion appear to support Ward's' value for the slope of the Log N versus I/'T curve for naphthalene. The somewhat lower slope given by ethylene chloride2 may possibly be explained on the grounds of slight solvation, as Hildebrand and JenksS account for the similar case of naphthalene in chloroform.
FIQ.z
TABLE VI Experimental Slope
Lr of Cio Hs
Calculated
960 970 975 975
4400 cal. 4440 4460 4460
Lower Limit of Straight Line Function jooc 60 50
65
summary The solubility of naphthalene in ethylene chloride, ethylene bromide, ethylidene chloride and ethylidene bromide has been determined. The slopes of the upper portion of the log N versus I / T curves have 2. been calculated: 3 . The data seems to indicate the value 4440 calories for the molal latent heat of fusion of naphthalene. I.
J. Phys. Chem., 30, 1316(1926). I 5 (1924))gives a log S versus I/Tcurve for naphthalene the curve is almost the same as ours, as nearly as can be jud e i f r o m the small figure. I t is stated that Ward has supplied the data but thus far we 8ave seen no published data on this system. a J. Am. Chem. Soc., 42, 2180 (1920).
* Hildebrand ("Solubility,"
in eth lene chloride; the slope
05
