SPECIFIC HEATS AND RELATED PROPERTIES OF T H E BINARY SYSTEM METHYL ALCOHOL-TOLUENE
L. S. MASON
AND
E. ROGER WASHBURN
Chemistry Laboratory, University of Nebraska, Lincoln, Nebraska Received December 13, 19S6
Departure of the behavior of binary liquid solutions from ideality has provoked a number of attempts to correlate measurements of physical constants with the state, and change of state, of molecular aggregation within the solution. The conclusions that may be drawn from a study of one type of constant, however, are often but partially sustained by, or wholly inconsistent with, those deduced from studies of some other constants. It is at present rather generally presumed that mutual solution of two or more liquids may be accompanied by dissociation of complex groups of molecules present in either or all of the component liquids into simpler units, and by aggregation of unlike molecules as a result of compound formation or the so-called process of solvation. Inconsistencies in the application of this view are attributed t o variation in the sequence and degree in which the association, dissociation, and solvation occur as a result of the different thermodynamic conditions in which measurements of different constants are performed. The vagueness of such assumptions emphasizes the qualitative nature of our knowledge about these processes. Various investigators have examined diverse phases of the problem of solution ( 2 , 3 , 4 , 5, 6, 7, 10, 18). In this laboratory property-composition data are being accumulated for binary systems of low molecular weight alcohols with low molecular weight hydrocarbons of the aromatic series and with six-carbon cycloparaffins, and for ternary systems of the above combinations with water as the third component (11, 14, 15, 16). It is hoped that an extensive study of such systems will reveal more of the mechanism of solution than is a t present known, and that results of such studies may afford generalizations more widely applicable than those in vogue. During the course of these investigations interest in the pertinence of heats of mixing to the problems of solution has arisen. Knowledge of the specific heats of solutions is a prerequisite of the determinations of heats of mixing, and a study of the literature reveals but very few data on the specific heats of solutions of organic liquids. The purpose of the present work was to devise a convenient and efficient technique for determining 481
482
L. 6. MASON AND E. ROGER WASHBURN
the specific heats of liquid mixtures and for measuring the heat changes attending the solution of the liquids, and to study these and other properties of the system methyl alcohol-toluene, MATERIALS
Methyl alcohol. The methyl alcohol used was a synthetically prepared product of a degree of purity originally greater than 99.5 per cent. This alcohol was desiccated over lime and carefully fractionated in an all-glass still. The relative density of the purified product was d;? = 0.78672, and the refractive index was nz" = 1.32659. Toluene. The best grade of toluene obtainable from the Eastman Kodak Company was treated repeatedly with freshly cut sodium and fractionated. No change in density was observed. A sample showed only a faint yellow coloration after standing over concentrated sulfuric acid for several hours. The relative density of the material was dtt0 = 0.86229 and its refractive index n;" = 1.49365. Toluene obtained from the Mallinckrodt Chemical Company was used for supplementary determinations of the thermal quantities after it had been treated with concentrated sulfuric acid, then sodium, and finally fractionated. EXPERIMENTAL
Apparatus and procedure Specific heats of solutions of methyl alcohol and toluene through the concentration range were measured adiabatically a t 25°C. and 35°C. by a method involving the use of a Dewar flask in a hand-controlled air bath as a calorimeter. Heat was supplied to the liquids by a heating coil of nichrome wire of measured resistance, which was connected to two 6-volt storage batteries and a silver coulometer in series. The magnitude of the current passed was calculated from the weight of silver deposited in the coulometer and the length of time of deposition, the latter being measured with a stopwatch. The temperature rise was measured on a Beckmann thermometer which passed through a stopper fitting tightly in the Dewar flask. A series of five copper-constantan thermocouples indicated when any difference of temperature existed between the liquid in the flask and the air of the bath, and the bath was heated or cooled to restore equality of temperature. The heat supplied for each determination was calculated from Joule's law. All solutions for these and other measurements were prepared with weight pipets. The heat capacity of the calorimeter, which was found to be 20.86 cal. per degree at 25°C. and 21.41 a t 35"C.,was determined by using toluene and methyl alcohol as "standard" liquids in the calorimeter. The values are the averages of groups of determipations in which the average deviation from the mean is less than 0.7 per cent. The values used for the specific heats of methyl alcohol and toluene are those of Bose (1) and of
BINARY SYSTEM METHYL ALCOHOL-TOLUENE
483
Williams and Daniels (17), respectively. The specific heats of methyl alcohol and toluene recalculated on the basis of the determined heat capacity of the calorimeter are 0.610 and 0.392 a t 25°C. and 0.613 and 0.402 a t 35°C. The calorimeter was a Dewar flask of approximately 200-cc. capacity fitted with a cork stopper coated with water glass and equipped with a Beckmann thermometer graduated in 0.005", a vertical glass stirrer, five thermocouples, and a nichrome heating coil. The series of thermal junctions was connected to a critically damped wall galvanometer whose deflections were observed with the customary arrangement of telescope and scale. The thermocouples were used only to indicate if a difference of temperature existed inside and outside the calorimeter. The resistance of the heating coil, which was found to be 7.730 ohms a t 25"C., wasmeasured potentiometrically by balancing the I R drop across the coil against the I R drop across a Bureau of Standards 10-ohm resistance. The temperature coefficient of the resistance of the heating coil exerted a negligible effect on the quantities measured a t 35°C. The air bath was a wooden cabinet 3 x 3 x 2 feet in dimensions, equipped with a window to permit observation of the interior, a large fan for circulating the air, and heating and cooling units which permitted rapid adjustment of the temperature of the interior to any desired value. A long metal shaft, attached eccentrically to a pulley above, passed through the roof of the cabinet and operated the stirrer in the calorimeter. The Beckmann thermometer in the calorimeter was read through the window of the cabinet by means of the telescope on a cathetometer. The silver coulometer consisted of two electrodes of sheet silver, approximately 13 x 10 x 1mm. in dimensions, supported by silver wires welded to them. These electrodes were suspended in a 1-1. beaker filled with an electrolyte, prepared according to the method of Wartenburg and Schutaa (12) and stirred by a rotary stirrer. The cathode was weighed, before and after each deposition, alter it had been thoroughly washed with distilled water, dried in an electric oven a t 125°C. for ten to fifteen minutes, and allowed to cool twenty minutes. The electrodes were interchanged after six depositions had been made on one of them. The silver coulometer proved to be convenient and reliable for measurements of current strength. A good quality stop-watch graduated in tenths of a second was used to measure the lengths of time that current passed in the circuit. Heat changes of mixing were measured for the concentration range a t 25°C. and 35"C., using the same calorimeter as for the measurements of specific heats. A small flask, also equipped with thermocouples, was suspended in the cabinet in such a manner that a weighed amount of one liquid could be delivered directly into the calorimeter to be mixed with a weighed amount of the other liquid. The entire system was closed and
484
12.
S. MASON AXD E. ROGER WASHBURN
in temperature equilibrium before mixing. Temperature changes occurred within one minute, and it was found unnecessary to maintain adiabaticity during such a short time interval. It was necessary, however, to determine the heat capacity of the calorimeter for this short interval of temperature change. The value 16.9 cal. per degree was used for calculations a t 25°C. and 17.6 at 35°C. The densities of the series of solutions were determined a t 25°C. and 35"C., using pycnometers of the type designed by Wade and Merriman (13). Volume changes were calculated from the measured densities. TABLE 1 Specific heats WEIGHT PEI< CENT -4LCOHOL
1
1
WEIGHT O F SOLUTION
TEMPERATURE RISE
T
= 25T.
3 3 3 4
10 17 18 71 28 37 37 43 52 27 60 77 70 94 80 34 89 55
84 83 83 82 81 80 79 79 79
440 743 025 678 452 944 991 493 041
0 0 0 0 0 0 0 0 0
2910 2974 2971 3065 3268 3310 3174 3028 3337
9 18 28 37 47 57 68 78 89
85 83 83 82 81 81 80 79 78
030 791 273 236 608 675 555 627 231
0 0 0 0 0 0 0 0 0
314: 3063 3149 3124 3128 3034 3068 3048 3081
T 17 73 09 96 66 63 21 47 25
1
CATHODE W'EIGHTS
4
4 4 3 3 =
1
TIME IN SECOND 8
1
SPECIFIC H E A T
920 996 856 083 021 058 095 873 985
540 340 540 540 600 600 540 5 10 600
0 0 0 0 0 0 0 0 0
453 474 503 518 547 564 581 600 607
104 044 086 833 970 951 126
588 576 564 552 540 528 516 504 492
0 0 0 0 0 0
526 541 560 577 597 610
35°C
'
4 4 4 3 3 3 4
Refractive indices of the solutions were measured with reference to the D line of sodium a t 25°C. and 35°C. with a Bausch and Lomb immersion type refractometer. RESULTS
Some of the values obtained for the specific heats of the methyl alcoholtoluene solutions a t 25°C. and 35°C. are given in table 1. The first column to the table, from left to right, expresses the weight per cent of alcohol in the solutions; the second, the weight of solution used in the determination;
485
BINARY SYSTEM METHYL ALCOHOL-TOLUENE
the third, the weight of silver deposited in the cathode of the silver coulometer; the fourth, the temperature rise observed on the Beckmannthermometer; the fifth, the time in seconds during which the determination was performed; and the sixth, the calculated specific heat of the solution.
LL u
%
o
io
20
30
40
50
70
60
eo
100
90
WEIGHT PER CENT bf ALCOHOL
FIQ.1
--
c
B
e
TABLE 2 Changes in specific heat and partial molal specific heats
--
a
si
siI
si
1
d
c
2
o"
1 : f
sgi
12
o"" 00.0 10.o m.0 30 .O 40 .O 50.0 60 .O 70.0 80 .O 90.0 100 .o
0.402 ).392 0,392 0.441 0.414 +8.2 0.454 0.48 0.436+10.3 0.486 0.50i 0.457+10.3 0.510 0.52' 0.479 +9.4 0.528 0.54 0.501 +8.0 0.545 0.55! 0.523 +6.9 0.562 0.571 0.545 +5.7 0.580 0,591 0.566 +4.6 0.597 0.602 0.588 +2.7 0.610 0.61( 0.610 0.613
--
rn
lo" - __
0.402 0.423 0.444 0.465 0.486 0.508 0.529 0.550 0.571 0.592 0.613
+7.3 +9.5 +9.7 +8.6 +7.3 +6.2 $5.5 +4.6 +3.0
36.1 34.7 32.2 29.7 27.6 25.7 24.2 22.9 21.8 20.7 19.5
37.0 35.2 32.6 30.1 27.8 25.9 24.4 23.1 22.0 20.9 19.6 ~
31.1 28.7 25.1 22.8 21.4 20.5 20.4 20.3 20.1
36.3 36.7 37.8 39 0 40.1 41.2 41.5 41.7 43.0
29.8 27.6 25.0 22.9 21.7 20.5 20.4 20.3 20.1
37.3 37.6 38.4 39.6 40.5 42.0 42.2 42.2 43.2
-
Duplicate determinations on solutions of the same concentration in all but a few cases vary less than 0.5 per cent. Figure 1is a graphical representation of the data shown in table 1. The upper curve shows the values obtained at 35OC.and the lower those a t 25°C.
486
L. 8. MASON AND E. ROGER WASHBURN
In table 2 the first column is the weight per cent of alcohol in the solutions; the second, the specific heat of the solution at 25°C. read from an enlarged graph like figure 1; the third, the specific heat of the solution
0
10
20
30 WEIGHT
&?
70
40
50
PFR
CENT OF ALCOHOL
60
90
100
FIQ.2
WEIGHT PER CENT AIEOAOL
1
w?l:J!
1
TABLE 3 Heats of m i x i n g c,
OF MIXTURE
I
TEMPERATURE CHANGE
1
H E A T OF MIXING I N CAL. PER GRAM
T = 25°C. 10 38 19 62 29 80 38.82 48.28 57.64 68.07 78.58 88.68
10 10 18 86 37 74 57 74 78 61
86.067 84 592 85.299 81.922 80.930 81.214 80.395 78,999 79.046
85 223
84 023 82 032 81 150 79 337
0.450 0.480 0.504 0,521 0.538 0 554 0.573 0.590 0,602
1, I ~
-1 -1 -0
626 882 a33 412 909 404 852 303 712
-2.343 -2.639 -2.692 -2.482 -2,174 -1,832 -1 450 -1.047 -0.581
-3 -3 -3 -2 -1
721 983 702 504 285
-2 -2 -2 -1
-3 -3 -3 -3 -2
-2
T = 35T. 0 455 0483 0 523 0 559 0 595
1
462 758 731 943 -1 050
BINARY SYSTEM METHYL ALCOHOL-TOLUENE
487
tion). The fifth, sixth, and seventh columns give the same values for 35°C. I n the eighth and ninth columns are listed the molal specific heats of the solutions at 25OC. and 35"C., calculated by dividing the observed specific heat at each concentration by the sum of the number of moles of each constituent, assuming 1 g. of solution in all cases. I n the last four columns are shown the partial molal specific heats, of alcohol and toluene at 25°C. and 35"C., obtained by a graphical method described by Lewis and Randall (8). I n figure 2 the deviations of specific heats from ideality are plotted against the weight per cent of alcohol in the solutions at 25OC. and 35"C., the upper curve being for 35°C.
e,,
FIQ.3
Heats of mixing expressed in calories per gram of solution are tabulated in table 3, and are plotted against weight per cent of alcohol in figure 3. The upper curve represents the values at 35°C. The specific heats of the solutions used in the calculations were obtained from the graphs of the specific heats versus composition at the two temperatures. In table 4 are found the observed densities and the volume contractions which attend mixing of the liquids at 25°C. and 35"C., together with the refractive indices and per cent deviations of refractive indices from a linear relationship for both temperatures. Volume changes are expressed in per cent of the volumes which the solutions would have had, had there been no change on mixing. Volume changes are plotted against concentration of alcohol in the solutions at both temperatures in figure 4. I n these curves the greater deviations are observed at 25°C. b
488
L. 8. MASON AND E. ROGER WASHBURN w
-0.160
$ v
-0120
%, $ 5 -0.080
0 0
IO
20
50
40
30
60
WEIGHT PER CENT
70
80
90
-100
OF ALCOHOL
FIQ.4
WEIGHT PER LENTALCOHOL
I
TABLE 4 Changes in volume and vefractive indices OBSERVED DENBITY
T 00 00 9 17 18 67 28 09 37 94 48 41 59 73 68 92 80 00 89 42 100 0
00 00 9 17 18 67 28 09 37 94 47 66 57 54 67 97 78 47 89 25 100 0
=
c~~~~~~~
1
WEIGHT PER CENTALCOEOL
85294 84571 83833 83116 82354 81617 80862 80087 0 79313 0 78518 0 77738
oBBERvED
~
1
CHANGE IN REFRACTIVE INDEX
RE::izE IN PER CENT T = 25'C.
25'C.
0 86229 0 85528 0 84799 0 84081 0 83330 0 82519 0 81666 0 80977 0 80156 0 79463 0 78672 ____ T = 35°C. 0 0 0 0 0 0 0 0
1
00 .oo 9.68 18.46 29.20 38.87 48.41 59.73 68.92 80.00 89,42 100 .o
000 -0.0~
-0.105 -0,139 -0.159 -0,148 -0,139 -0.126 -0.100 -0.069 000
-
1.49365 1.47622 1,45903 1.44120 1.42523 1 ,40873 1,39045 1 37544 1.35756 1.34289 1.32659
000 -0.085 -0.259 -0.254 -0.244 -0.286 -0.245 -0.221 -0.180 -0.103 000
T = 35°C.
106 112 122 105 097 081 -0 042 000 -0 -0 -0 -0 -0 -0
I
28 09 37 94 47 66 57 54 67 97 78 47 89 25 100 0
1.48839 1.47110 1.45416 1.43824 1.42222 1 .40527* 1,38952 1.37249 1.35533 1.33907 1.3240
000 -0.151 -0.242 -0.280 -0.273 -0.333 -0.308 -0.310 -0.300 -0.192 000
DISCUSSION
If the heat added to solutions of methyl alcohol and toluene were utilized only for raising the temperature of the solutions, specific heat would be a straight-line function of composition, with the line terminating in the
BINARY SYBTEM METHYL ALCOHOL-TOLUENE
489
specific heats of the pure components. The specific heats of the solutions, however, are greater than those calculated from additivity for all concentrations, as is shown by figure 1. Figure 2 emphasizes the extent of the deviations and shows the concentrations at which the deviations are most pronounced. Toluene is regarded as a normal liquid and methyl alcohol as strongly associated. The heating of solutions of the two liquids during specific heat determinations probably causes partial dissociation of the methyl alcohol, and this most markedly in the solutions of low alcoholic content. It appears that the presence of toluene enhances this dissociation by a dilution effect. The large heat capacities of the solutions of low alcohol concentrations are due, in a measure at least, to the absorption of heat involved in the process of dissociation of the alcohol aggregates. The plotted data indicate that the changes of specific heat from ideality are greater at 35°C. than at 25"C., which would suggest that a greater dissociation occurs at the higher temperature. The partial molal heat capacities, (values listed in table 2), are greater for all values calculated than the molal heat capacities of the pure liquids. Williams and Daniels (18) attribute such behavior to the formation of a compound of the components, or to the fact that heat capacity is increased by an increase in the number of molecules even though the total weight remains unchanged. It would seem that the latter effect is predominant in the system under consideration in view of the large negative heats of mixing, which furnish evidence that compound formation is unlikely. The difference between the partial molal heat capacity in solution and the molal heat capacity of the pure alcohol is greatest in the mixtures of low alcohol content, which would indicate a greater number of molecules and hence greater positive deviations from ideality for the heat capacities of the mixtures in that region of concentration. The view that dissociation of the alcohol is aided by the dilution with toluene is supported by a consideration of the heat changes which attend solution of the liquids. At all concentrations, and a t both 25OC. and 35"C., heat is absorbed on mixing, and the greatest absorptions are in the solutions of high concentrations of toluene, and these are greatest at 35°C. Evidence, in the case both of heat capacity deviations and of heats of mixing, indicates that more dissociation occurs at the higher temperature and at the lower concentrations of alcohol. It would be expected, however, that measurements at successively higher temperatures would ultimately show smaller and smaller deviations and smaller heat changes as the alcohol became more and more dissociated. Large negative heats of mixing indicate the absence of compound formation or solvation in that these processes are usually associated with an evolution of heat. However, it is possible that such processes occur to a small extent, or with small evolution of heat, or conceivably, with a nega-
e,,
490
L. 8. MASON AND E. ROGER WASHBURN
tive heat of formation. Under these conditions their thermal effects would be obscured by the heat absorption of dissociation. It would be expected that an increase in the number of molecules in a solution, by virtue of dissociation, would be accompanied by an increase in volume, but in the case of the present system apparent dissociation of the alcohol is accompanied by a contraction in volume, and the largest contractions are found in approximately the same concentration areas as are the greatest deviations of specific heats and the greatest absorptions of heat on mixing. The evidence adduced from the thermai quantities indicates that dissociation is increased a t the higher temperature, but volume contractions are found to be decreased. If a combination of alcohol and toluene molecules occurs in the manner suggested in connection with heats of mixing, the volume contractions might be reconciled. Such a combination, however, would be assumed to be quite unstable, and undoubtedly would be broken up when the temperature of the solutions was raised during determinations of heat capacities. Washburn and Lightbody (15) observed that mixtures of ethyl alcohol with both benzene and toluene display both contractions and expansions of volume on mixing, and that methyl alcohol and benzene show both expansion and contraction. They postulated that a change from volume expansion to volume contraction is due to a shifting of the equilibrium among simple, complex, and compound molecules with changing concentration (9), but why methyl alcohol and toluene should be unique among these four systems in showing a contraction throughout the concentration range it is difficult to say, especially in view of the fact that the similarity of the liquids would lead to the expectation that the mixtures would behave similarly. Measurements of volume changes a t other temperatures should be of help in answering this problem. Refractive indices vary slightly from a straight-line function of composition, and the refractive indices of the mixtures are lower than the values obtained from calculations on the basis of an ideal solution. A decrease of refractive index is contrary to what would be expected if the only molecular change occurring in the solutions is one of dissociation of the alcohol. The supposed increased number of molecules together with a contraction in volume should produce an increase in the refractive indices. Here, as in the case of volume changes, a union of alcohol and toluene molecules may be the governing factor in causing the decrease in refractive indices. It would seem, however, that influences other than relative numbers of molecules are in effect, because the maximum deviations of refractive indices occur in solutions of greater alcohol content than do the greatest contractions in volume. Some changes in the chemical nature of the molecules, not yet understood, must be responsible for the situations observed. As yet it is possible to make only qualitative interpretations of the type of
BINARY SYSTEM METHYL ALCOHOL-TOLUENE
49 1
data presented here, but more complete knowledge of this and similar systems should furnish evidence for roughly quantitative predictions. SUMMARY
A rapid and efficient technique for measuring specific heats of solutions of organic liquids has been developed. The apparatus used, with slight modification, is readily applicable to the determination of heat changes attending solution of the liquids in each other. Specific heats and heats of mixing of methyl alcohol-toluene solutions have been determined at 25OC. and 35°C. for the entire range of concentration. The heat capacities observed for the solutions are greater than those calculated from additivity for all concentrations. Mixing of the liquids is accompanied by an absorption of heat in all cases. Both effects are most pronounced in concentrations of low alcohol content. The changes in volume resulting from mutual solution of the liquids, and the refractive indices of the solution, have been measured throughout the concentration range at 25°C. and 35OC. Solution of the liquids is accompanied by a contraction in volume; refractive indices are lower than those calculated on the basis of additivity. The thermal effects observed are accounted for by assuming that methyl alcohol, which in the pure state is associated, undergoes dissociation into simpler molecular aggregates during mixing of the liquids and during subsequent heating in the determinations of the heat capacities of the solutions. A union of the constituent molecules has been assumed to explain the observed changes of volume and refractive index. REFERENCES
(1) B O ~ EE.: , 2. physik. Chem. 68, 585 (1907). International Critical Tables, Vol. V, p. 114. (2) BRAMLEY, A. : J. Chem. SOC.109,496 (1916). (3) Bossy AND BUIGNET:Ann. chim. phys. 4 , 5 (1865). R. B.: Trans. Faraday SOC.8 , 2 0 (1912). (4) DENISON, F.: 2. physik. Chem. 64, 727 (1908). (5) DOLEZALEK, (6) HARTUNG, E. J.: Trans. Faraday SOC. 12,66 (1916). (7) K E ~ E SD., B., AND HILDEBRAND, J. H. : J. Am. Chem. SOC.39,2126 (1917). M.: Thermodynamics and the Free Energy of (8) LEWIS, G. N., AND RANDALL, Chemical Substances, p. 38. The McGraw-Hill Book Co., New York (1923). (9) MADGIN, W. T., PEEL,J. B., AND BRISCOE,H. V. A.: J. Chem. SOC.1927,2873. (10) PARKS, G. S., AND KELLEY,K. K. : J. Phys. Chem. 29,727 (1925). E. R.: J. Am. Chem. SOC.64,4217 (1932). (11) VOLD,R., AND WASHBURN, (12) VON WARTENBURO, H., AND SCHUTZA, E . : z. Elektrochem. 36,254 (1930). (13) WADE,J., AND MERRIMAN, R. W. : J. Chem. SOC.101,2429 (1912). E. R., HNIZDA,V., AND VOLD,R. : J. Am. Chem. SOC.63,3237 (1931). (14) WASHBURN, E . R., AND LIOHTBODY, A.: J. Phys. Chem. 34,2701 (1930). (15) WASHBURN, (16) WASHBURN, E. R., AND SPENCER, H. C.: J. Am. Chem. SOC.66,361 (1934). J. W., AND DANIELS,F.: J. Am. Chem. SOC.46,903 (1924). (17) WILLIAMS, J. W., AND DANIELS,F.: J. Am. Chem. SOC.47, 1490 (1925). (18) WILLIAMS,
