SOME PHYSICAL-CHEAIlCAIi PROPERTIES OF MIXTURES OF ETHYL A N D ISO-PROPYL ALCOHOLS BY GEORGE S. PARKS AND K E N N E T H K. KELLEY
In a recent study' of some physical-chemical properties of mixtures of ethyl and normal propyl alcohols, the I;iystem was found to approximate to the requirements of an ideal or "perfect" solution to a remarkable degree. Thus the process of forming the solutions was accompanied by a very small heat absorption, in no case more than five calories per mol of resulting mixture, and by a volume shrinkage averaging only .cz 5%. Moreover, the partial pressures of each component as calculated from the experimental data a t 25' C were but slightly higher than those called for on the basis of Raoult's law. Indeed, the data obtained were so striking that the (question a t once arose as to whether a study of other binary systems involving closely related alcohols would yield analogous results. Accordingly the two 3ystems, ethyl and iso-propyl alcohols, and iso-propyl and n-propyl alcohols, we re investigated in an essentially similar manner. The present paper presents the data obtained in the study of mixtures of the former pair of liquids. In constitution iso-propyl alcohol differs from ethyl alcohol merely by the substitution of a methyl group for one of the hydrogen atoms in the carbinol group of the latter compound. Hence in several important physical-chemical properties the two alcohols are very similar, as the following table indicates.
TABLE I CH3CIJpOH
Dielectiic Constant at zoo C Capillary Constant a t 3'I C4 Association Factor a t 3 I' C4 Relative Internal Pressures5 a t zo°C From surface tension (yc V") From critical data 5 ( IO5)(U f
v2)
From coefficients of expansion and compressibility 5 T
(P )
IC-~
25.8' I .c8 2.74
(CH3)lCHOH
26(f1)~ I
.Oj
2.86
3.4
-_
4.9
3.5
2.6
2.8
2.8
Parks and Schwenck: J. Phys. Chem. 28, 720 (1924) 2. physik. Chem. 29, 24G (1899). Lowe: Wied. Ann. 66, 398 (1895). 4Ramsay and Shields: 2. physik. Chem. 12, 468 (1893). These relative internal pressures were calculated by the methods described by Hildebrand: J. Am. Chem. SOC.,41, 1067 (1919). The necessary data were obtained fromqhe Tabellen of Landolt-Bornstein-Roth ( 1 9 1 2 ) .
* Abegg and Seitz:
728
GEORGE S. PA4RKSAND KENKETH K. KELLEY
The equality in dielectric constants and the relatively small differences in the capillary constants and degrees of association indicate that we are dealing with two liquids of the same degree of polarity. A comparison of the values for the relative internal pressures is not quite so satisfactory, since the three different pairs of results obtained by the various methods of calculating internal pressures are not very concordant. On the average, the values for iso-propyl alcohol run about 14CI, lower than those for ethyl alcohol. This difference, altho appreciable, is hardly great enough t o lead us to expect any large deviations from the laws of a perfect solution; since, as Hildebrand has remarked (1oc.cit.) “a very decided difference in internal pressures is required to produce any considerable deviation from Raoult’s law.” On the other hand, the two preceding factors-chemical similarity and equal polarity-point definitely in the other direction and would lead us to suspect that this pair of liquids might form a series of solutions which would be almost “perfect.”
Experimental PuriJication of Substances. A good grade of 95% ethyl alcohol was dehydrated by a preliminary distillation over lime in the ordinary manner, followed by a second distillation over a small quantity of calcium metal. The resulting product was carefully fractionated and the middle portion, about 60% of the total, was selected for use in the present investigation D:f 0.78516, which corresponds to 99.97% ethyl alcohol according to the U. S. Bureau of Standards tables’. The iso-propyl alcohol was prepared ia exactly the same way. After the fractionation process the final product had a density of 0.78086 2 s o / 4 O , which corresponds to 99.99yc alcohol on the basis of Brunel’s value2 of 0.78084 for xooyc, and the variation per I qcof water of 0.00230 obtained by Lebo3. Hence it is safe to conclude that our purified materials had much less than .I% water at the start cf the following series of measurements. Formation of the alcohol mirclurcs. The alcohols thus prepared were used in making two series of five mixtures each, which varied systematically in steps of approximately 16yG of each component. Thus the first mixture in series A contained by weight I 5.95% of iso-propyl alcohol and 84.05Y0 of the ethyl alcohol; the second mixture, 33.26y0 iso-propyl and 66.74yGethyl, etc. About I O O cc of each mixture was made up. In the course of the preparation of these solutions the heat of mixing was determined. In this measurement the apparatus and procedure used in the determination of the heats of mixing of ethyl and n-propyl alcohol were again employed. For the details concerning these the reader is referred to the earlier article4. U. S.Bureau of Standards: Scientific Paper, KO.197 (1913). J. Am. Chem. SOC.45, 1336 (1923).
* Brunel: 4
Lebo: J. Am.Chem. SOC.43, 1006 (1921) Park? and Schwenck: loc. cit.
PROPERTIES O F MIXTURES OF ALCOHOLS
729
The data obtained in the present investigation appear in columns four and five of the following table. In contrast with the results found for the ethyln. propyl alcohol system, the process of forming the various solutions in this instance took place with the evolution of heat. The numerical values of the heat effect, while not great, are about two and a half times those obtained in the earlier study-an indication that the solutions in the present case may not be as “perfect.”
TABLE I1 Heat of Formation of the Mixtures a t Liquid
% CiHsOH by weight
I 2A
2B 3A 3B 4A 4R
5A
SB 6A 6B *
Mol fraction of CiHhOH
100.00
I .ooo
84.05 8 3 . 19 66.74 47.05 49.92 49.83 33.67 33.49 16.95
.873 .866
17.72 0.00
.219
.724
,726 .j65 ’ 565 ,398 ,396 .210
.000
25’
€€eat of mixing in calories per mol of per gram of mixture mixture __ +.I52
+.I49
+.204
+ +.a42 +.238 +
,201
.212
+.
207
+.I75
+. 168 ___
+ 7.3 + 7.2
+IO. 2
+ro.o s 1 2 . 7
+12.4 JrII.6 +II.3 S I O . 0
+9.5
The numbers given in the first column to the various liquids are for convenience in reference in subsequent pages of this paper. Thus, when we mention “liquid No. 4A” for instance, we shall be referring to the solution containing 49.92% (by weight) of ethyl alcohol. Densities and Total Pressures. The densities of the liquids were next determined in the usual manner, a double-walled, evacuated specific gravity bottle of I O cc capacity being used for this purpose. All weighings were corrected for the buoyancy of the air, and the final values appear in the second column of Table 111. For purposes of comparison determinations were also run on samples of the pure alcohols after they had been put thru the mixing apparatus; this was done because it seemed inevitable that the samples would absorb traces of moisture during such a procedure and we desired to have all our liquids in as comparable a condition as possible. On the assumption that n-e are here dealing with perfect solutions, the densities of the mixtures were calculated by use of the relationship I
D
-
PI 100 d,
+- P:,d2 100
where dl and dz are the densities of the components in the pure state, PI and PZ are their corresponding weight percentages in the resulting solution and D is
73 0
GEORGE S. PARKS AXD KENNETH K. KELLEI’
the density of the solution. The values, thus obtained, appear in column three of the following table; they average only . o I ~ lower , than the experimental results. In fact, we are practically justified in concluding that, within the limits of experimental error, the volume of the solution is the sum of the volumes of the two pure components-a characteristic relationship of the ideal solution.
TABLE I11 (Temperrtture 25’ C) Liquid
Density Observed Calculated
57.2
2B
0.7851
0.7851 0.7851
3A 3B
0.7845 0.7846
0.7844 0.7844
4.A
0.7839 0.7840
0.7839 0.7839
0.7833 0.7834 0.7829 0.7829 0.7820
0.7833 0.7832 0.7827 0.7827
SA 5B 6A 6B
7
Ideal
59. o mm (assumed)
2A
4B
Vapor Pressure
_-
0.7857 0.7852
I
0bserv ed
57.2 54.9 55.3 52.4 52.6 49.7 49.5 47.6 47.5 44.4
J J
57.2 mm 57.0
55.0 55.0 52.5 52.6 50.2
50.2 47.5 47.6
In column four are tabulated the vapor pressures of the various liquids, determined by the differential method described in the previous paper. These results have been obtained on the assumption of 59.0 mm. as the vapor pressure of pure ethyl alcohol and are probably reliable to + 0 . j mm. For purposes of comparison the “ideal” vapor pressures have been calculated by Raoult’s law. Throughout the entire range of concentrations the two sets of values, observed and ideal, agree within the limits of experimental error.
Refractive Indices. At this stage of the investigation the two series of solutions, A and B, were combined. The new series of liquids, thus produced, had the compositions indicated in Table 117. The refractive indices of the various liquids were then determined with a Zeiss-Pulfrich refractometer. These measured values agree remarkably well with those calculated (column j) on the basis of the assumption of a straightline relationship between the index of refraction and the composition, by weight, of the solution. Thus this measurement provides an easy and rapid method of analyzing an unknown mixture of these alcohols. As the instrument used could be read with a percision of f I minute and the refractive angles for the two pure alcohols differ by 168 minutes, the accuracy of the method is about 0.6%~
PROPERTIES O F MIXTURES O F ALCOHOLS
731
TABLE IV (Temperature % CJlsOH
Liquid
by weight
25'
Mol fraction of CzHsOH
C) Refractive Index Observed Calculated
2
83.61
1.000 .869
1.3595 I . 3622
I . 3620
3
66.89
.725
I . 3646
I . 3646
4
49 87 33.60
565 .393
1,3673 1.3697
1.3673 I . 3698
17.23
.214
I
.3721 1.3750
--
100.00~~
I
9
5 6 7
0.00
'
,000
1.3723
Vapor Composition ut 25" C. The composition of the vapor phase in equilibrium with the solutions at 2 5 " C. was nest determined. This was accomplished by passing air (freed from mater and carbon dioxide) thru a series of three bubblers, each containing about 2 0 cc of the mixture under consideration. The air thus saturated with the vapor of a mixture was then passed thru a condensing tube immersed in liquid air; the alcohol separated out as a solid glass on the walls of this tube and, when about I cc of distillate had been collected, was analyzed by measurement of its refractive index. Two separate determinations were made on each mixture; the results appear in Table V. TABLE V Ethyl Alcohol in the Vapor of the Liquids a.t 25' C Liquid
Toby weight
I
IO0
2
8 7 . 5 and 8 7 . 5 72.9 " 7 4 . 6
3 4
5 6
7
594 41.8 20.4
,) 'I
,)
Mol Fraction (mean value)
I . 000
59.0 40.6 22.5
0
. yo1 785 .654 .479 .263
'
. 000
Viscosities. We also determined the viscosities of the various liquids, using an Oswald viscosimeter in a 25" C thermostat, regulated to .oI'. The value 0.00893 dynes per sq. cm., as obtained by Hosking', was assumed for the water which was used in standardizing the instrument. The time, measured by a stop-watch, averaged around I O O seconds and thus limited the accuracy of these results t o .t.00003 dynes per sq. cm. Hosking: Proc. Roy. Soc. S. 8. Wales, 43, 37 (1909).
GEORGE S. PARKS AND KENNETH K. KELLEP
732
TABLE VI Viscosities at Liquid
I 2
3 4
5 6 7
Observed values
,01080 ,01182 ,01299 .or415 .01607 . 01796 ,02048
2 j”
C (in dynes per sq. cm.) Calciilated [Kendall’s equation)
. 01 I84 .01306 ,01450 ,01614 .or806
Calculated (Logarithmic equation)
.orI74 . or288 .or426 .OI. 588 .OI786
Comparison of the experimental results with the data calculated by Xendall’s cube-root equation1 (7% =xlq1x+x2q4$2,where 71 and qsare the viscosities of the pure components and X1 and X2 are their respective mol fractions) show the latter to be too high by an average value of 0.4%. On the other hand the logarithmic relationship, log q =XI log q1 X2log 72, gives results which are on the average ,0.9% low. Evidently Kendall’s equation, while not entirely satisfactory, is the better approximation for ideal solutions Qf this type.
+
Summary Reviewing the results of the various measurements, we find that (I) A small heat evolution (in a,ll cases less than 13 calories per mol of resulting mixture) takes place on formation of the several solutions. (2) An extremely small volume shrinkage, on the average only .oI%, accompanies the process. (3) The measured vapor pressures of the resulting liquids agree within the limits of experimental error with the “ideal” values calculated by means of Raoult’s law. (4) The retractive indices of the solutions are practically a straight line function of their weight compositions. ( 5 ) The observed viscosities for the various solutions exhibit on the average a negative deviation of 0 . 4 7 ~from Kendall’s cube-root equation. The only appreciable deviation from the laws of the perfect solution is found in the heat effect on formation of the mixtures and even in this case the departure from the ideal is relatively small. Judging the data as a whole, we are led to the conclusion that the system under discussion is almost “perfect.” Furthermore, in view of the fact that a similar situation was found t o exist in the case of mixtures of ethyl and n-propyl alcohol, it seems probable that this conclusron will be found valid for all systems involving closely related alcohols. Department of Chemzstry Stnn.ford Unicersaty October 28, 198.4.
Kendall: J. Am. Chem. SOC. 42, 1776 (1920).
