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 con- cerning 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 - .02 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 require- ments of Kentiall’s “cube-root” equation,2 the 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. More- over, 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 factors- chemical dissimilarity, cliff erent 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

9 Kendall: J. Am. Chem. Soc., 42, 1776 (1920).

(1925); Winchester: hlaster’s Thksis, Stanford University (1923).

410 GEORGE S. PARKS A S D CLARE 6. CHAFFEE

TABLE I Dielectric Capillary Association Relative Internal Constant' Constant? Factor? Pressure3

(C'H 3 ) ?C 0 2 1 . j 1.82 1.26 I , 3 2 ICH,)?CHOH 26 I . 0 j 2 . 8 6 2 ' 19

Experimental Purijcatz'on of Subs tances . C. P. materials served as the starting point in

preparing the acetone and iso-propyl alcohol for the present investigation. To 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. I ts 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-Rijrn- stein "Tabellen."

The iso-propyl alcohol \vas first dehydrated by two distillations over lime in the ordinary manner and was then carefully fractionated. The final pro- duct had a density of 0 . ~ 8 1 3 0 z ~ O , ' ~ O : which corresponds to 99.8';T alcohol on the hasis of Rrunel's \due4 of 0.;8084 for 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 com- positions 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 determin- ation the proper weight of the component which was to be present in greater quantity was 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 tightly- fitting 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

1 Landolt-Bbrnstein-110th-Scheel: "Tabellen," pp. 1036 1037 (1923). 2 Ramsay and Shields: Z. physik. Chem., 12, 468 (189.3). 3 Mortimer iJ. .\m. ('hem. Foc., 45, 640, 1923), gives ~ . , p and 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 th 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 ~ 6 11921).

MIXTURES O F ACETONE AXD ISO-PROPYL ALCOHOL 441

such a way that it could be kept nearly full. This container occupied a posi- tion 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 con- t 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 a t 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 wait- ing interval and again reading the thermome- tlsr, 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 con- sideration. 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, respec- tively. 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 represen- ted graphicaIly in Fig. 2 . The process of f'srming the various solutions took place with the ahsorption of considerable heat-an indica-

FIQ. I The Calorimeter for the Deter- mination of the Heats of Mixing.

tion 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 to a temperature lowering of 11.7' .

Densitzes and Refractive Ind ices . 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

0 .2 .4 .6 .8 1.0' tometer, the method of Moore'for tem- MOLE FRACTION perature measurement being employed.

FIQ. 2

Fraction of Iao-Propyl Alcohol.

These observed results fall on a curve Heat of Mixing plotted against Mole WThich runs below that of the values

calculated on the basis of a straight- line 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 iso- propyl 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 182 minutes, the accuracy of the method is about 0.57~.

The composition of the vapor phase in equilibrium with the solution a t 25' C was next determined. This was ac- complished 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 alcohol- acetone 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 measure- ment of its refractive index. The results appear in the third column of Table IV.

Vapor Composition at 25" C.

1 Moore: J. Phys. Chem., 25, 281 (1921).

where dl and d2 are the densities of t he components in the pure state, PI and Pz are their corresponding weight percent- ages in the resulting solution and D is the density of the solution. Using this equation, we have calculated the densi- ties 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 refrac-

MIXTURES O F .4CETOSE AND ISO-PROPYL ALCOHOL 443

Liquid

Liquid

TABLE I1 Heat of Formation of the Mixtures a t 20' C.

1so:Propyl Alcohol Heat of mixing in calories c; by weight Mole fraction per mole of mixture

0 . 0 0 ,000 - 1j.65 . I j2

16 .56 ,161 - 190 - 196

3 0 . 0 0 ,292 -316 30.89 ,302 - 320 46 .34 ,452 -362 49 .72 ,486 -388 4 9 . 8 5 ,488 -385 66 .90 ,661 -343 6 7 . 4 1 ,667 -35; 67.61 ,669 -338 81, 70 ,812 -236 85.37 ,849 -197 _- I 0 0 , O O 1.000

T.4BLE 111 (Temperature, 25' C)

Dmsity Observed Calculated

0.7855 0.7832 0 ,7831 0 .7818 0.7816 0 . 7 8 0 5 0 .7806 0 . 7 8 0 7 0.;801

--

0 . j 848 0.7848 0 . 7 8 4 2 0 .7842 0.7836 0.7834 0 . 7834 0 . :82;

0 . j 8 0 4 0 . 7 8 2 0 0 . 7 8 0 5 0.7819 0 ,7813 ~-

Refractive Index Observed Calculated

1 .3555

1.3579 I . 3604 1 ,3605

I ,3641 I , 3641 I ,3672

1 . 3 j j 8

I ,3634

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 a t 25' C were also meas- ured 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. Dur- ing 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 . CHAFFEE

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.

Liquid

TABLE I V Vapor Composition a t 25' C

In the original liquid 1101 Fraction of Iso-Propyl .ilcohol

In the vapor . I j2 . 108 ,161 ,090

, 2 9 2 ' 13; ,302 , 1 3 7 ,452 . I ; 2

,486 , 2 0 2

. 488 19; ,661 , 2 6 5

,812 ' 3 9 0 ' 8 4 9 , 4 2 2

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 to a resi- due of dissolved air in the liquid or possibly to a slight change in composition of the liquid during the pumping process. In the cases of pure acetone and pure iso-propyl alcohol these errors were entirely eliminated, since by a pro- cess 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 jo

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.

Liquid

I

2B 3c 4B 5A 6C 7

TABLE I' Vapor Pressures at 25' C.

Mole Fraction of Total Vapor Pressure Iso-Propyl Alcohol Observed Ideal

,000 226. j mm ,161 221.6 ' ' .33I 190.0 166.2 ,486 16;,2 138.0 ,661 139.6 106. I

197.2 mm

, 8 2 5

I . 000

IO0,O 76. I

44 ' 3

1 hlathews: J. Am. Chem. SOC. 48, 574 (1925). Parks and Kelley: J. Phys. Chem., 29, 730 (1925).

MIXTURES O F ACETOSE 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 solu- tions mere calculated, These results appear in Columns 2 and 4 of the follow- ing table. The "ideal" pressures in all cases were derived on the assumption of Raoult's law:

PA = NAP0.4 where PA and XAi are 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 f pressures are represented graphically

z in Fig. 3. It is obvious that the solu- tions show a marked positive deviation - from Raoult,'s lax-, amounting in the case W of an equimolal mixture to about 21%. E

3 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 to .oI'. The a time was measured by a stopwatch. The value 0.00893 dynes per sq. cm., MOLE FRACT~JN as obtained by Hosking,' was assumed for the water which was used in standar- Total and Partial Pressures plotted

dizing the instruments. Propyl Alcohol.

T'iscosities.

0 .2 .4 .e Ip

FIG. 3

against the Mole Fraction of Iso-

TABLE 1'1 Partial Pressures a t 2 j" C

Liquid Partial Pressure of .Icetone

Experimental Ideal Experimental Ideal

Partial Pressure of Iso- Propyl Ilcohol

I 226 . jmm __ o . o m m ~

3 c 1 6 2 . 5 I jI. j " 2 7 . 5 1 4 . j "

4B 134.2 1 1 6 . 5 3 3 . 0 2 1 . j

j -1 102.6 7 6 . 8 3 7 . 0 29 .3 6C 59 ' 9 3 9 . 6 40. I 3 6 . 5

2R 199, j ' ' 190. I mm 2 1 . 9 ' I j . x m m

__ 4 1 . 3 __ i 0.0

Comparison of the experimental results x i th the data calculated by Xendall's cube-root equation? (7% = xlql% + -u27?%, where 71 and 7 2 are the

* Hosking: Proc. Roy. SOC. S. S. Wales, 43, 37 ( 1 9 ~ 9 ) . * Kendall: J. Am Chem. SOC., 42, I 7 7 6 (1920).

446 GEORGE S. PARKS .4SD CLARE S. CHAFFEE

viscosities of the pure components and s1 and 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.”

TABLE VI1 T-iscosities a t 2 5 ’ C (in dynes per sq. cm.)

Liquid Observed values Calculated values , 0 0 3 0 8

2 -1 ,00347 ,00446

2u ,00349 3 A , 0 0 3 8 ~ ,00607

I ___

___ 3B ,00395 4A ,00486 ,00833 4B

j A , 00703 , 0 1 2 o j

6A. .010;1

. o o j o i ~- 4c . O O j I 3

___ 613 ,0118; , 0 1 6 2 2

I , 0 2 0 2 0

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 34yc from 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.” In this connection it is interesting to note that the comhination of deviations from the behavior of a perfect solu- tion-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-

Hildebrand: “Polutiility,” p., 63 (1921,). \Ye wish to take advantage of the present 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 OF 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 iso- propyl alcohol molecules tend largely to 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 cube- root 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 Uniterst1 y , Calzfornza. October 27. 1926.