SOME PHYSICAL-CHEMICAL PROPERTIES OF MIXTURES OF ACETONE Ah’D ISO-PROPYL ALCOHOL
BY GEORGE S. PARKS AND CLARE S. CHAFFEE
A few years ago, as a preliminary t,o a st,udy of the equilibrium between pcetone, hydrogen and iso-propyl alcohol, some information was needed concerning the physical-chemical properties of liquid mixtures of acetone and iso-propyl alcohol. In the absence of any such datata,the investigation to be described in the present paper was undertaken. This particular study is the more interesting because of the recent investigations’ in this laboratory of the three binary systems formed from ethyl, normal propyl, and iso-propyl alcohols. These three binary systems were found to approximate to the requirements of an ideal or “perfect” solution to a remarkable degree. Thus the process of forming the various solutions was accompanied by a very small heat effect, the maximum AH values being f4.8, -12.7 and - 10.1 calories per mole of resulting mixture in the respective cases of the ethyl-n-propyl alcohol system, the ethyl-iso-propyl alcohol system and the n-propyl-iso-propyl alcohol system. The corresponding volume changes averaged - . 0 2 jc;C, - ,0107~ and - .008c:. The measured vapor pressures for the three systems agreed within the limits of experimental error with t’he ideal values calculated by Raoult‘s law. The viscosities approximated in magnitude to the requirements of Kentiall’s “cube-root” equation,2the respective deviations averaging only +.iyC, -.4C; and f.75. All these results, of course, are not especially surprising, inasmuch as the three liquids involved are closely related alcohols of approximately the same polarity and internal pressure. Hon-ever, in the pair of substances under consideration in the present, paper we have a somewhat different situation. One is an alcohol and the other a ketone, the second being the dehydrogenation product of the first. Moreover, while both liquids might he classed as moderately polar, & study of the values in Table I for dielectric constant, capillary constant, and association factor indicates that the alcohol is decidedly the more polar and the more abnormal of the two. As far as relative internal pressures are concerned, the difference is even more striking. Hence, in view of these important factorschemical dissimilarity, clifferent degrees of polarity, and inequality of internal pressures-it might reasonably be predicted that this binary system would exhibit’ marked departures f r o m the behavior of the perfect solution. And such was found to he the case. 1 Parks and Schwenck: J. Phvs. Chem., 28, 720 (1924); Parks and Kelley: 29, 727 (1925); Winchester: hlaster’s Thksis, Stanford University (1923). 9 Kendall: J. Am. Chem. Soc., 42, 1776 (1920).
410
GEORGE S. PARKS A S D CLARE 6. CHAFFEE
TABLE I Dielectric Constant'
(C'H3 ) ?C0 ICH,)?CHOH
Capillary Constant?
Association Factor?
21. j
1.82
1.26
26
I.0j
2.86
Relative Internal Pressure3 I ,32 2
'
19
Experimental Purijcatz'on of Substances. C. P. materials served as the starting point in preparing the acetone and iso-propyl alcohol for the present investigation. T o remove possible traces of met,hanol and water, the C.P. acetone wa% treated with anhydrous calcium chloride at the rate of 2 0 0 gm. per liter. The resulting mixture was allowed to stand for several clays: with occasional shaking, and was then distilled. The product thus obtained was furt,her treated with a small quantity of anhydrous calcium chloride and carefully fractionated. The middle portion. representing about jocp of the total antl boiling hetween 56.10' and .;6.24', was selected for use in the measurements. Its denait,y was o.;8j j at 2:' 4', a result, which compares favorably with the two values for 1 0 0 ~ acetone, ; 0.;863 and 0,7849,given in the Landolt-Rijrnstein "Tabellen." The iso-propyl alcohol \vas first dehydrated by two distillations over lime in the ordinary manner and was then carefully fractionated. The final product had a density of 0 . ~ 8 1 3 0z ~ O , ' ~ O : which corresponds to 99.8';T alcohol on the hasis of Rrunel's \due4 of 0.;8084for roo";( and the variation per 1'; of water of 0.00230 o h i n e d by I ~ b o . ~ Detemz'ncition of Heat of -1Jixz'ng. The pure liquids thus prepared were used in making a number of mixtures. which had the designations and compositions represented in C'olumns I , 2 antl 3 of Table 11. In the course of the prcpnration of these solutions the heat of mixing was determined a t an average teniperatnre of 20' C'. The method employed for this measurement was somewhat different, from that described by Parks and Schwenck and should give more accurate results. In any particular determination the proper weight of the component which was to be present in greater quantity w a s placed in a 400 cc Dewar jar, Fig. I , equipped with a cslit~w,tetl thermometer accurate to .oI', a stirrer of the propeller type and a tightlyfitting cork cover to exclude moisture. The required weight of the othcr component was then placed in a srnall copper container, A , of about 60 cc capacity, which had a bent outlet tube extending upward from its hottoni in Landolt-Bbrnstein-110th-Scheel: "Tabellen," pp. 1036 1037 (1923). Ramsay and Shields: Z. physik. Chem., 12, 468 (189.3). 3 Mortimer iJ. .\m. ('hem. Foc., 45, 640, 1923), gives ~ . ,and p 2.90 for the respective internal pressures of acetone and ethyl alcohol, relative to that of naphthalene as standard. Since Parks and Iielley found that the internal pressure of iso-propyl alcohol was, as the average of several different modes of calculation, about 115; below that of ethyl alcohol, we have accordingly reduced t h r value 2.90 by 1 4 ~ and ; in this way have obtained 2 . 1 for ~ iso-propyl alcohol. Brunel: J. Am. C'hem. Soc., 45, 1336 ( 1 9 2 3 ) . jLebo: J. Am. Chem. Soc., 43, 1 0 ~ 611921). 1 2
MIXTURES O F ACETONE AXD ISO-PROPYL ALCOHOL
441
such a way that it could be kept nearly full. This container occupied a position in the lower part of the Dewar jar and thus the two liquids to be mixed were brought to the same initial temperature. When this condition had heen reached, the contents of the copper vessel were forced over into the Dewar jai outside by a stream of air entering the cont tiiner via the upper entry tnbe. The air used i'or this purpose had been previously saturated with acetone or alcohol, as the case might be, in order to avoid any cooling effect due to vaporization within the calorimeter system. The resulting mixture was immediately stirred for two minutes at a rate of 70 R. P. M.and t:hen allowed to stand for a similar period after which the final temperature was taken. By at once repeating the stirring and waiting interval and again reading the thermometlsr, the temperature correction due to stirring, heat exchanges with the surroundings, ew., was obt,ained. The corrected temperature change was then used for calculating the heat effect in producing one mole of the solution under consideration. This calculation was made in .the usual way, by using 16 calories as the heat capacity of the calorimeter and the values of €'arks and Kelley' for the specific heats of pure acetone and iso-propyl alcohol, respectively. The heat capacity of each mixture was taken as the sum of the heat capacities of the pure components, a procedure which is strictly accurate only in case there is no deviation from the laws of the perfect solution. The results, which are good to ~7~ or kletter, are given in Table I1 and are represenFIQ.I ted graphicaIly in Fig. 2 . The process of f'srming the various solutions took place with The Calorimeter for the Determination of the Heats of Mixing. the ahsorption of considerable heat-an indication that we are dealing with solutions that are far from perfect. In the case of the equimolal mixture this heat absorption reached a maximum value of 387 calories, equivalent t o a temperature lowering of 1 1 . 7 ' .
Densitzes and Refractive Indices. The densities of the liquids were next determined in the usual manner, a specific gravity bottle of 35 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. 1
Parks and Kelley: J. Am. Chem. SOC., 47,
2 9 1
(1925).
442
GEORGE 6. PARKS
AND CLARE 6. CHhFFEE
In the formation of a perfect solution the resulting volume should be equal to the sum of the original volumes of the components involved, or in terms of densities the relationship is
where dl and d2 are the densities of t h e components in the pure state, PI and Pz are their corresponding weight percentages in the resulting solution and D is the density of the solution. Using this equation, we have calculated the densities which appear in Column 3 of the table. These values average about 0.3% greater than the observed densities of Column 2 ; in other words, there is actually a very appreciable volume in crease accompanying the formation of these solutions. The refractive indices (Table 111, Column 4) for sodium light were then determined with a Zeiss-Pulfrich refrac0 .2 .4 .6 .8 1.0' tometer, the method of Moore'for temMOLE FRACTION perature measurement being employed. FIQ.2 These observed results fall on a curve Heat of Mixing plotted against Mole WThich runs below that of the values Fraction of Iao-Propyl Alcohol.
calculated on the basis of a straightline relationship between the index of refraction and the composition by weight of the solution. Measurement of the index of refraction provides an easy and rapid method of analyzing an unknown mixture of acetone and isopropyl alcohol. As the instrument used could be read with a precision of *I minute and the refractive angles for the two pure substances differ by 1 8 2 minutes, the accuracy of the method is about 0.57~.
Vapor Composition at 25" C. The composition of the vapor phase in equilibrium with the solution a t 25' C was next determined. This was accomplished by passing air (freed from water and carbon dioxide) through a series of three bubblers, each containing about 20 cc of the mixture under consideration. The air thus saturated with the vapor of a mixture was then passed through a condensing tube immersed in liquid air; the alcoholacetone vapor 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. The results appear in the third column of Table IV. 1
Moore: J. Phys. Chem., 25, 281 (1921).
MIXTURES O F .4CETOSE AND ISO-PROPYL ALCOHOL
443
TABLE I1 Heat of Formation of the Mixtures a t
20'
C.
Heat of mixing in calories 1so:Propyl Alcohol c; by weight Mole fraction per mole of mixture
Liquid
0.00
,000
1j.65 16.56 30.00 30.89 46.34
. I j2
- 190
49.85 66.90 67.41 67.61 8 1 , 70 85.37
,161 ,292 ,302 ,452 ,486 ,488 ,661 ,667 ,669 ,812 ,849
-316 - 320 -362 -388 -385 -343 -35; -338 -236 -197
I00,OO
1.000
49.72
T.4BLE
- 196
_-
111
(Temperature, 25' C ) Liquid
Observed
0.7855
Dmsity Calculated --
0.7832 0,7831
0 . j848
0.7818
0.7842 0.7842
0.7816
0.7805
0.7848
0.7807
0.7836 0.7834 0 . 7834
0.;801
0 . :82;
0.j804
0.7820 0.7819
0.7806
0.7805 0,7813
Refractive Index Observed Calculated
1.3555 1.3jj8 1.3579 I . 3604 1,3605 I ,3634 I ,3641 I , 3641 I ,3672
1.3584 I ,3586 1.3611 I , 3613 I . 3642 1.3649 1.3649 I ,3681
~-
T-apor Pressures at 25" C. The vapor pressures at 25' C were also measured in the case of several representative liquids. The method employed for this purpose was as follows. In each case a 5 cc sample of the liquid under consideration was introduced into a small bulb which was connected to a mercury manometer. The air in and above this liquid was removed by a pump consisting of activated charcoal immersed in liquid air and the bulb and manometer system were then sealed off from the outside atmosphere. During the pumping process the liquid in the bulb was protected from vaporization by its immersion in a bath of liquid air, a t the temperature of which its vapor
444
GEORGE S . PARKS A S D CLARE S . C H A F F E E
pressure was negligible. The bulb and contents were then hrought to a temperature of 25' C and the vapor pressure of the liquid was measured on the manometer by means of a cathetometer.
TABLE IV Vapor Composition a t
25'
C
1101 Fraction of Iso-Propyl .ilcohol I n the original liquid In the vapor
Liquid
. I j2
. 108
,161
,090
,292
'
,302 ,452
,137 . I ;2
,486
,202
. 488 ,661
19; ,265
,812
'390
'849
,422
13;
The results appear in Table V, Column 3 . With the mixtures the results in some cases may possibly involve errors as great as 2 or 3pc, owing t o a residue of dissolved air in the liquid or possibly to a slight change in composition of the liquid during the pumping process. I n the cases of pure acetone and pure iso-propyl alcohol these errors were entirely eliminated, since by a process of partial evaporation of the liquid all dissolved air could be pumped off without the risk of a change in composition. Our result for acetone at 2 j o agrees well with the value 226.3,recently obtained by h1athew-b' in a very careful investigation. For iso-propyl alcohol our present result is almost identical with that of an earlier measurement*in this laboratory.
TABLE I' Vapor Pressures at Liquid
25'
Mole Fraction of Iso-Propyl Alcohol
C. Total Vapor Pressure Observed Ideal
226.j mm 221.6''
I
,000
2B 3c 4B
,161 .33I ,486
5A
,661
16;,2 139.6
,825
IO0,O
6C 7 1
I . 000
hlathews: J. Am. Chem. SOC. 48, 574 (1925). Parks and Kelley: J. Phys. Chem., 29, 730 (1925).
190.0
44 ' 3
197.2mm 166.2 138.0 106.I 76.I
M I X T U R E S O F A C E T O S E AXD ISO-PROPYL ALCOHOL
445
For each component in a given mixture the partial pressure is simply the product of its mole fraction in the vapor phase and the total pressure of the solution. Accordingly, from t'he experimental dat.a in Tables I V and V the partial pressures of the acetone and iso-propyl alcohol in the various solutions mere calculated, These results appear in Columns 2 and 4 of the following table. The "ideal" pressures in all cases were derived on the assumption of Raoult's law: PA = NAP0.4 where PA and XAiare respectively the partial pressure and the mol fraction of component X in a given solution and Po, is its vapor pressure in the pure state. The data for the partial and total pressures are represented graphically in Fig. 3. It is obvious that the solutions show a marked positive deviation
f
z-
from Raoult,'s lax-, amounting in the case W of an equimolal mixture to about 21%. E
3
T'iscosities. We also determined the fl viscosities of the various liquids, using a W two Oswald viscosimeters in a z j o C (1 t,hermostat, regulated t o .oI'. The a time was measured by a stopwatch. 0 .2 .4 .e Ip MOLE FRACT~JN The value 0.00893 dynes per sq. cm., as obtained by Hosking,' was assumed FIG.3 for the water which was used in standar- Total and Partial Pressures plotted against the Mole Fraction of Propyl Alcohol.
dizing the instruments.
Iso-
TABLE 1'1 Partial Pressures at Liquid
2 j"
C Partial Pressure of IsoPropyl Ilcohol Experimental Ideal
Partial Pressure of .Icetone Experimental
Ideal
2R
226.jmm 199, j ''
__ 190. I mm
21.9
3c 4B
162.5 134.2
I jI.j
27.5
14.j 21. j
j-1
102.6
33.0 37.0
I
6C
59 ' 9
i
0.0
o.omm
"
116.5 76.8 39.6
40.I 41.3
__
'I
~
j.xmm "
29.3 36.5
__
Comparison of the experimental results x i t h the data calculated by -u27?%,where 71 and 7 2 are the Xendall's cube-root equation? (7% = xlql%
+
* Hosking: Proc.
* Kendall: J. Am
Roy. SOC. S . S.Wales, 43, 37 ( 1 9 ~ 9 ) . Chem. SOC.,42, I 7 7 6 (1920).
446
GEORGE S. PARKS .4SD CLARE S. CHAFFEE
viscosities of the pure components and s1and s2 are their respective mole fractions) shows that the calculated values are too high on the average by 34%. As Kendall’s equation has been found to give results within 0 . 7 7 of the experimental data in the case of the three alcohol mixtures previously studied, it may be considered as one of the properties of a perfect solution. On the basis of this test the present system is far from “perfect.”
T-iscosities a t Liquid I 2
-1
2u
3A
3B 4A 4B 4c
TABLE VI1 C (in dynes per sq. cm.)
25’
Observed values ,00308
,00347 ,00349 ,0038~ ,00395 ,00486
Calculated values
___ ,00446
,00607
___
,00833
.oojoi .OOjI3
~-
jA
,
6A.
.010;1
___
613
,0118;
,01622
I
,02020
00703
,012oj
Summary and Conclusion Reviewing the results of the various measurements, we find that (I) A considerable heat absorpt,ion, amounting to 387 calories per mole in the case of the equimolal mixture, takes place on formation of the several solutions. (2) An appreciable volume increase, on the average 0 . 3 ~ $ ,accompanies the process. (3) The measured vapor pressures and partial pressures of the resulting liquids considerably exceed the ideal values calculated by means of Raoult’s law. (4) The observed viscosities for the various solutions exhibit on the average a negative deviation of 34ycfrom Kendall’s cube-root equation. Judging the data as a whole, we are led to the conclusion that the system under discussion is far from “perfect.” I n this connection it is interesting to note that the comhination of deviations from the behavior of a perfect solution-heat absorption and volume increase on mising, together with a positive deviation froin Raoiilt’s lalv-is entirely consistent with the findings in studies on other syskms.’ For an explanation of the present results, we have re\Ye wish to take advantage of the present Hildebrand: “Polutiility,” p., 63 (1921,). opportunity to express our appreclatlon of Professor Hildebrand’s excrllent monograph. I n preparing this paper we have freely used the ideas expressed therein and are indebted particularly to Chapter Y l l , “Causes of Deviations from Raoult’s Law.”
MIXTURES O F ACETONE AND IYO-PROPYL ALCOHOL
447
course to the great differences in association factors and relative internal pressures of acetone and iso-propyl alcohol, as shown in Table I. These differences indicate that the molecules of iso-propyl alcohol have a much greater attraction for one another than for acetone molecules or than the acetone molecules have for each other. In the pure state, therefore, the isopropyl alcohol molecules tend largely t o associate, although such association is undoubtedly indefinite and not strictly stoichiometric. On the formation of solutions with acetone, these attractive forces in the iso-propyl alcohol are weakened and as a result there is a dissociation or break-up of the more or less indefinite associated groups or “group” molecules, accompanied by an absorption of heat and an increase in volume. This decreased association as far as the alcohol is concerned leads to its positive deviation from Raoult’s law in the series of solutions; while on the other hand the inherent tendency of the alcohol molecules to associate, though weakened, produces a squeezing out effect on the acetone and positive deviations for this component also. The large negative deviations of the observed viscosities from Kendall’s cuberoot equation can also be attributed to the lessened forces of attraction and decreased association of the iso-propyl alcohol molecules in the various solutions. Department of Chemistry, Stanjord Uniterst1y , Calzfornza. October 27. 1926.
