T H E CHANGES I K VOLUME AND TEMPERATURE THAT ACCOMPANY T H E M I X I S G OF ORGANIC LIQUIDS. I* BY E. ROGER TVASHBURN AND ALBERT LIGHTBODY
Few American chemists seem to have studied the changes in volume and in temperature which accompany the mixing of organic liquids, although several European scientists have studied these phenomena. Bussy and Buignet' measured the changes in temperature caused by the formation of several binary mixtures. They were especially interested in alcohol-carbon disulfide and alcohol-ether, and with these systems, they also measured the change in volume. Guthrie,? another pioneer in this work studied the changes shown by alcohol-carbon disulfide and triethylaminewater mixtures. Clark3 measured these changes for alcohol-water, propyl alcohol-water, and aniline-xylene systems of various concentrations. He also measured the heat of mixing for several other solutions. Patterson and Montgomeriei studied mixtures of ethyl tartrate in various solvents such as water, ethylene bromide, nitrobenzene, methyl alcohol, benzaldehyde, quinoline. Among other workers who investigated isolated systems, measuring one or both of the changes mentioned, are Young,; Werriman,6and Holmes.' I n more recent times we find the work of van Laar,&dealing with the theoretical aspect of heat of mixing as related to other physical constants of liquids. Perrakisg studied the changes in volume and in temperature as shown by the following mixtures; acetic acid-benzene, o-cresol-alcohol, and each of the four lower alcohols with benzene. His work will be mentioned later in this paper in connection with some of our own experiments. Madgin and Briscoelo made a broad survey of the field, measuring the temperature changes shown in 628 equal volume, binary, mixtures chosen from 37 liquids. In later work" they present careful studies of both temperature and volume changes for a few of these systems. Even with these researches and others as a background, no satisfactory theory relating the phenomena connected with the formation of solutions of organic liquids has been advanced. Some of the systems show a contraction in volume accompanying a rise in temperature; in others it accompanies a
* Contribution from the Chemistry Laboratory of the University of Kebraska. 'Ann. Chim. Phys., (4) 4, 5 (186j). ? Phil. Mag., ( j ) , 18, 495 (1884). Physik. Z., 6 , 154 (1905). J. Chem. SOC.,95, 1128 (1909). J. Chem. SOC.,83, 4 j (1903). J. Chem. S O C ,103, 1774 (1913). J. Chem. SOC.,103, 2147 (1913). * Proc. Acad. Sci. Amsterdam, 35, 399 (1923). 'Compt. rend., 178, 703, 1482 (1924); J. Chim. phys. 22, 164, 280 (192j). 'OJ. SOC.Chem. Ind., 46, I O ~ (1927). T "Briscoe, Peel, and Madgin: J. Chem. SOC., 1927, 2873; J. Phys. Chem., 32, 2 8 5 (1928).
2702
E. ROGER WASHBURN AND ALBERT LIGHTBODY
fall in temperature. Other systems show an expansion in volume to be assocciated with absorption of heat, while in still others it occurs with an evolution of heat. To make the situation still more complex some systems show a contraction in volume when one of the components predominates and an expansion when the other is in excess. Obviously there are many exceptions to the generalization’ that, usually when heat is given off during the mixingof liquids, a contraction in volume, and a negative deviation from Raoult’s law also occurs. Holmes (loc. cit.) believes that physical forces alone are operative in solutions of this kind; that the size of molecules, intermolecular attraction (of the nature of gravitation), friction, etc., are the factors which determine the changes in physical properties which accompany the mixing of liquids. Bramley2 presents some evidence to show that compounds are formed when certain organic liquids are mixed. Madgin, Peel, and Briscoe (loc. cit.) are inclined to believe that association or combination of molecules is the important factor. Investigators have been handicapped in their attempts to formulate an adequate theory explaining these phenomena by a scarcity of reliable data. The early work has either been very general in its nature, covering a wide range of materials with a limited range of concentration, or it has involved the studies of unrelated liquids. It has occurred to the authors that a systematic study of mixtures, covering the entire range of concentration, of chemically related materials might lead to a more satisfactory generalization concerning these solutions than has yet been advanced. For our initial step in this effort, we have chosen to work with mixtures of the lower alcohols, with benzene, and with toluene. Apparatus The calorimeters were constructed usingunsilvered Dewar flasks, fitted with a cork stopper, thermometer, stirrer, and a small opening through which the tip of a pipette might be inserted. The stirrer, a glass rod bent in the form of a circle a t the bottom was worked up and down around the thermometer by means of a motor, which gave it approximately forty-eight strokes per minute. The calorimeters were operated in a constant temperature bath at 25.0 f. r o c . during all determinations. The thermometer used for the smaller temperature changes was a five-degree Beckmann while the one used for the larger temperature changes was a tenth-degree thermometer having a larger range. The accuracy of the thermometers was checked by comparison with a thermometer calibrated by the Bureau of Standards. The pipettes used in introducing all the liquids into the calorimeter were carefully selected and calibrated. They ranged from 5 cc. to 50 cc. I n order to obtain some concentrations, two or more pipettes were used in introducing the first liquid into the flask, but never more than one was used in adding the second liquid to the first. The time of flow from even the largest pipette 1 2
J. H. Hlldebrand: “Solubility,” pp. 61-62 (1924). J. Chem. SOC., 109, IO, 434, 469, 498 (1916).
VOLUME-TEMPERATURE CHAiiGCS FOR ORGASIC SOLUTIOSS
2 703
was about forty seconds. They were all allowed to drain carefully and the last drop mas forced out of the tip because the difference in surface tension of the liquids used caused drops of different sizes to be retained. I n order t o determine the change in volume n-e measured the density of the resulting solution and compared it with the density which the solution would have had if there had been no volume change. The density measurements were made using calibrated, glass stoppered pycnometers of the type used by Wade and Nerriman.' The volumes of the pycnometers ranged from thirty-three to thirty-five cubic centimeters, so that an average error of one milligram in each of the three weighings necessary for a density determination would cause an error of only twelve parts in one hundred thousand in the measured density. All weights were corrected to vacuo. Materials The benzene and toluene were of the highest quality obtainable from Nallinckrodt Chemical Company. Their purity was established by comparison of the boiling points, densities, and refractive indices lyith the values given in the International Critical Tables. The methyl alcohol was Merck's C. P. quality. Tests showed it to be very nearly pure. The ethyl alcohol was a standard absolute alcohol. I t was dried with metallic calcium and fractionated in an all glass apparatus. Iso-propyl alcohol was obtained from Eastman Kodak Company. The density mas a little too high, and attempts were made to dry it with metallic calcium. These resulted in no marked improvement so the untreated material was used. As the densities of different samples of these liquids varied slightly, they were determined before each run. The values for the liquids used are given in the tables of results. Method The calorimeter and the liquids to be used were first placed in the bath and allowed to come to temperature equilibrium. Then one of the liquids, usually the one used in the larger amount, was run into the calorimeter by means of a pipette. The temperature changes were then read every minute for at least five minutes before the addition of the second liquid and then for five minutes after this addition. I n each case the combined volumes equalled very nearly 100 cc. The maximum temperature change a a s usually reached in less than one minute. The values thus obtained were plotted against time and the actual temperature change obtained by erecting a perpendicular half the distance between the time of addition of the second liquid and the time of the greatest observed temperature change, extrapolating the curves until they cut the perpendicular and recording the difference so found. The stirring was continued for fifteen to twenty minutes to insure a homogeneous solution and then a portion was removed and transferred to the Wade and Merriman
Chem. SOC.,101, 2429
(1912)
2
704
E . ROGER WASHBURN A S D ALBERT LIGHTBODY
pycnometer which had previously been cleaned and dried. The transfer was accomplished by slowly exhausting the air from the bulb through a fine capillary tube as a steady stream of liquid went into the instrument. The pycnometer was then placed in the bath for thirty minutes when the level of the solution was lowered to the mark etched on the stem. After five minutes, to insure a correct adjustment of the meniscus, the pycnometer was removed from the bath, wiped with a slightly moist cloth and allowed to attain the temperature of the balance before weighing. At least eight or nine concentrations were studied with each pair of liquids, in order to cover the complete range in concentration. When there was some doubt as to the validity of a value, the determination was repeated. The value recorded in such cases is the average of the several values obtained. The water equivalents of the calorimeters were obtained by measuring the heat of neutralization of sodium hydroxide with hydrochloric acid. The concentration of the hydrochloric acid was established by the gravimetric determination of silver chloride. The sodium hydroxide was titrated with the acid. For the calculation of the heat evolved we used as the heat of neutralization a t Z ~ O C . ,13,640 calories, and as the specific heat of the approximately N/4 NaCl solution formed, .98 I calories.’ The results listed in table A are the average of several determinations made with each calorimeter. I n no case was the deviation from this average greater than ~ 0 . 0 1 calories. TABLE A cc
cc
Temp. c$nge C.
%&lN
8%”
51.80
51.93
3 .os
51.80
51.93
3 .oo
Heat developed calories
Heat used calories
Heat lost calories
Water equiv. degree
312.7
20.1
6.50
307.6
25.2
8.40
Calorimeter A 332.8
Calorimeter B 332.8
Results
I n the following tables we list the concentration of each solution used, expressed in mole fraction of alcohol; the observed and calculated density; the change in volume expressed in per cent of the calculated volume; the calorimeter used for the determination; the temperature change and the approximate heat change in calories per mole of mixture. The heat change is an approximation because the specific heats of only one or two of these mixtures are available, and these are of somewhat doubtful value because of the abnormally high density of the alcohol with which they were determined. The approximate heat changes were calculated on the assumption that the specific heat of the mixture is equal to the sum of the specific heats of each component multiplied by the mole fraction of the component in the mixture. Daniels, Mathews and Williams: “Experimental Physical Chemistry,” 79 (1929).
VOLCME-TEMPERATURE CHANGES FOR ORGANIC SOLUTIONS
2 705
Perrakis (loc. cit.) has shown that this assumption is not justified. We have included sufficient data so that when the specific heats are available the heats of mixing may be calculated. Curves showing how changes in volume and temperature vary with concentration are included.
TABLE I Methyl Benzene and Benzene Mole fraction alcohol
Observed density gms./cc.
Calculated density gms./cc.
Change in volume
Calorimeter
Change in temp.
Approx. heat change
0.000
0.8728j
-
-
-
-
0.104
0.86833
0.86856
0.03
A
3.43
127
0.196
0.86394
0.86426
0.04
B
4.09
151
0.279
0.8j961
0.8j99;
0.04
B
4.12
146
0.423
0.85118
0.8j139
0.03
B
4.26
140
0.68;
0.83014
0.82992
-0.02
B
3.32
94
0,868
0.80884
0.8084j
-0.05
0.926
0.80022
0.~9986 -0.03
0.9j2 0.97;
0.79571 o 79127
0.795jj 0.~9127
B A A A
0.34
I ,000
0.;8698
-
-
-
Calorimeter
Change in temp.
Approx. heat change
-
-0.02
-0.00
-
-
1.63
41
0.98
23
o,;j
1;
8
TABLE I1 Ethyl Alcohol and Benzene Mole fraction alcohol
Observed density grns./cc.
Calculated density gms./cc.
Change in
volume
-
0 000
0.8jzj6
0.074
0 . 8 6 ; ;6
0.86839
0.0;
A
3.14
I20
0 .r
0.8633j
0.86403
0.08
B
4.21
16j
0.8j880 o .8j062
0.85967
0.10
4.59
178
0.85093
0.82922
0.82911
-0.01
0.820
0.80jj4
0.80j28
-0.06
B R A A
0.896
0'79883
o.;9jjq
-0.16
4j
0.212
0.338 0.604
0.04
0.932
0.794j7
0.79418
-0.05
0.96;
0.79008
0.78982
-0.03
1.000
0.78545
4.73
182
3.90
140
2 .OI
70
.A
1.1;
40
X A
0 . jj
26
0.44
'5
2 706
E. ROGER WASHBUR;"; AND ALBERT LIGHTBODY
TABLEI11 Isopropyl Alcohol and Benzene Mole fraction alcohol 0.000
0.059 0.115 0.172
0.281 0.540 0.779 0.869 0.913 0.957 I ,000
Observed density gms./cc. 0.87285
Calculated density gms./cc.
0.86709 0.86201 0.86720 0.84797 0.82646 0.80661 0.79898 0.79513
0.86858 0.86430 0.86003 0.85150 0.83014 0.80878
0.17
X
0.27
B
0.33
€3
o 80024
0.16
R R B B
0.7959; 0.79169 -
0.11
A
I .87
0.09
h
-
-
0.94 -
-
0,79110 0.78 j 4 2
rr.tRLE
Change 1n volume
Calorimeter
-
-
0.42
0.45 0.2;
Change in temp.
Approx. heat change
-
-
3.56 4.93 5,72 6.66 6.75
139
4.23 2
.;I
204
239 284 308 198 '29
87 44 -
IT
;\[ethyl Alcohol and Toluene Mole fraction alcohol
Observed densitv gms./&.
0.000
0.86082
-
0.8j 8 2 0 0.85421
A B R
0.120
0.226 0.3'7 0.467 0.724
0.887 0.937 0.959 0.980 1.000
0.85061 0.84355 0.82534 0.8062I 0.798jg 0.79477 0.79089 0.78j06
Calculated densitv gms. /&.
85;53 85344 o 84975 o 84238 o 82394 o 80550 o j9813
Change in volume
Calorimeter
0
-0
0
-0
08 09
-0
IO
-0
B B
-0
13 17 09
-0
05 05
h A .A
-0
0
79444
-0
0
i9Oi.5
-0 02
x
-
Change in temp.
Approx. heat change
-
-
3.25
138 I57
3.70 3 96 3.89
I59 141
2
.8S
86
I
68 96
42
0
0.70
0.35 -
23
16 8 -
TABLE V Ethyl Ilcohol and Toluenp Mole fraction alcohol
Observed density gms./cc.
0.088
0.86079 0.8564;
0.169
0.85295
0.244
0,8493I
0.378 0.646
0.81225
0.845
0.911 0.942
0.80542 0.79764 0.79375
1.000
0
0.000
o ,82407
'
78558
Calculated density gms./cc.
Change in volume
Change in temp.
Approx. heat rhange
_-
-
-
A A B B
2.9i
4.07
I34 169 183
4.08
176
-0.11
B
3.02
-0.13
A A -1 -
I.24
119 44
0.88
31
Sj io3
0.07
0.8532; 0.84951 0.84200 0.82319 0.80439 0.79686 0.i9310
0.04
0 .
Calorimeter
0.02 -0
03
-0.10
-0.08
3.83
0.55
-
'9 -
VOLUME-TEMPERATURE CHANGES FOR ORGANIC SOLUTIONS
2fo7
3 Mo/e f r c c h o n a/Coho/ /n b e n z e n e
FIG.I
TABLE T I Isopropyl Alcohol and Toluene Mole fraction alcohol 0 000
o 069
135 0 I99 o 318 0 584 o 808 o 888 o 926 0
0
964
I 000
Observed densit) gms /cc. o 86072
Calculated density grns.1cc. -
-
85599 0 85174 0 84777 o 84006 o 82188 o 80469
0.8j712 0.85345
0.13
rl
0 . 2 0
B B B B B
0
0 0
0
79780 79434 79074
o 78742
Change in volume
0.84978
0.24
0.84246 0.82411 0.80j;i 0 ' 79843
0.29 0.27
0.14
Calorimeter
-
0.79476
0.0;
A A
0.79109
0.02
-1
-
-
0.08
-
Change in temp. -
3.23 4.36 j .OS
5.75 5.50 3 .59 2 .35 I .60
0.81 -
Appros. heat change
-
I49 210
243 277 271
I74 I11
76 38 -
2708
E. ROGER WASHBURN AND ALBERT LIGHTBODY
? FIG.2
Our results for ethyl alcohol and benzene mixtures do not agree with those of Perrakas, (loc. cit.). He obtained a contraction in volume only for one point at the extreme alcohol end of the curve (see dotted curve, Fig. 3) whereas we observe a contraction throughout a larger range in the concentration. From the density of the alcohol which he used it would appear that it contained about 3 or 4 per cent water, whereas our ethyl alcohol was 99.87% alcohol. We made up several alcohol-water mixtures of this concentration (97%) and on mixing them with benzene obtained values approximating those of Perrakas. We then made a study of 9 0 % alcohol mixed with benzene in different proportions and obtained the results tabulated in Table 7 and plotted in Fig. 3 . Evidently water has a pronounced effect on these changes. (The values for the dry alcohol-benzene mixtures are replotted for convenience of comparison.)
2 709
VOLUME-TEMPERATCRE CHANGES FOR ORGANIC SOLUTIONS
QO I
I
I
I
TABLE VI1 90% Ethyl Alcohol and Benzene fraction alcohol
Observed density gms./cc.
Calculated density gms./cc.
0.000
0 . 8 j 140
0.15
0.28
0.86298 0.85573
0.34
0.85227
0.40 0.61 0.82 0.93
0.84875 0.83548 0.81932 0.80996 0.80350
0.86461 0 . 8 j782 0.85442 0.85103 0.83755 0.82047 0.81039
Mo!e
1.00
-
Change in
volume
Change
in
volume
0.24
4.71 5.76
0.25
6.00
0.27 0.25
6.07 5.29
0.14
3.27
0.04 -
1.40
0.02
-
2710
E . ROGER WASHBURS A S D ALBERT LIGHTBODY
Discussion S o simple relationship exists between the temperature changes and volume changes for these solutions. Similar data for other mixtures must be obtained before important generalizations can be expected. I t will be noted that as we ascend the homologous series of alcohols with benzene and with toluene, the maximum temperature changes become greater for mixtures of corresponding concentration. The volume changes also seem to vary in a regular manner. Preliminary investigations with the fourth and fifth members of the alcohol series, however, seem to indicate that they would show smaller changes than the systems containing isopropyl alcohol. Madgin, Peel and Briscoe (loc. tit.) advance the idea that the following equilibria may exist in systcms of this sort. d, F t x h 7=?zAnBm R,FtyB Where A, and XX represent associated molecules and simple molecules in pure liquid A, By and yB in a pure liquid B. zAnBm represents a compound molecule which may be formed when A and B are mixed. If the dissociations to the simple molecules are endothermic and if the volume depends upon the relative number of simple, complex and compound molecules, then it may be possible to explain the various changes which we have observed on the basis of shifting the equilibrium of the different reactions. Water, a solvent favoring dissociation, if present in the ethyl alcohol-benzene system causes a larger drop in the temperature than is observed with dry liquids and a greater tendency for expansion throughout the concentration range. Of course the presence of the water makes a three component system and greatly increases the number of possible equilibria which may exist. The idea of polymerization of alcohol in benzene is not without supporting evidence from other sources. Freezing point studies of solutions containing more than very small amounts of alcohol seem to indicate that the alcghol is polymerized in the benzene. On the other hand, 8. F. Taylor,' found that the ratio of alcohol distributed between benzene and water indicated that alcohol has the same molecular form i n each liquid. It may be that the presence of the water causes the polymerized alcohol molecules to dissociate. Summary The changes in temperature and in volume which take place during mixing for the complete range of concentration with six pairs of organic liquids, have been presented. The effect of moisture upon the changes, as shown by one of the systems has been demonstrated; this shows the necessity of using dry liquids in this work.
t
J. Phys. Chem., 1, 463 (1896).
