SOME PHYSICAL-CHEMICAL PROPERTIES O F MIXTURES O F ETHYL AND N-PROPYTA ALCOHOLS BY GEORGE S. FARKS AKD JULIUS RAE SCHWEXCK
In recent years the concept of an ideal or “perfect solution’’ has become prevalent. Such a solution is one whose properties may he expressed in terms of the fractions of the components and the original properties of the pure substances by means of simple laws. Thus, in an ideal mixture of two liquids, the components should be miscible with one another in all proportions; the partial pressure of each component should follow Raoult’s law’; and the volume and heat content of the resulting mixture a t any temperature should be the sum of the volumes and the heat contents of the pure substances used -in other words, there should be no volume change or heat effect on formation of the mixture. From a priori considerations it would seem that mixtures of ethyl alcohol and normal propyl alcohol should closely approximate to these requirements. The two components involved are neighboring members of a homologous series and due t o this close chemical rclationship are very similar in certain important physical properties, as the following table will indicate.
TABLE I Dielectric Constmt2
CZHbOH CIH~OH
25.8 22.2
Cnl:illary Awocintioti Constant2 Facto12 2.74 I ,083
1.234
2.25
Internal Pressure4 Winther Traube
160
2030
2
1900
I800
The relatively small differences in the dielectric constants and the comparable values for the capillary constants and association factors indicate that the two liquids are about equally polar. Also the internal pressures, calculated by Winther from optical properties and by Traubc froim surface tension and van der Waal’s “a” and “b”, are of the same order of magnitude. All these factors-chemical similarity, equal polarity, and comparable internal pressures -would lead us to suspect that this pair of liquids niight form solutions which would be almost “perfect”. Up t o the present time, however, only very scanty experimental evidence on the subject has been available. Bose5 has determined the heat effect at 2 1 O c for the formation of solutions containing 49.6y0 and 32.9yo (by weight) of ethyl alcohol and found the values- 125 C R and ~ . - ,140cal. respectively C. N. 1,ewis: J. Am. Chrm. Soc. 30, 668 (19081. rlbegg and Seit,z: Z.physik. Chem. 29, 242 (1899!. 3 Ramsav and Phiclds: Z. physik. Chem. 12, 468 (,18931. Hildehrantl: J. Am. Chem. SOP.38, 1459 (19161. Rose: Cottingen Gcs. Xachr. 19C6, 333.
PROPERTIES O F MIXTURES O F ALCOHOL
72 I
for the heat of mixing per g a m of solution. These small values were almost within the limits of his experimental error and in a second publication' he briefly dismisses the subject with the statement that the heat of mixing of the alcohols studied (Le. methyl, ethyl, and n-propyl) was approximately zero. Hecent specific heat determinations2 on ethyl alcohol, n-propyl alcohol, and an rqiii-molal mixture of the two indicated that, within the limits of experimental error, the pure substances retain their original heat capacities after formation of the equi-molal mixture-a property of perfect solutions. These two researches seem to provide the only information applicable to this interesting binary system and it was for this reason that the work described in the present paper was undertaken. While the results obtained are of only moderate accuracy, it is felt. that the absence of any similar data renders their publication desirable.
Experimental Puri'cation of Substances. A good commercial grade of 95% ethyl alcohol was dehydrated by a preliniiiiary distillation over lime in the ordinary manner, followed by a second distillation over a small quant>ityof calcium metal. The result8ingproduct was carefully fractionat'ed and the middle portion, about 60% of t8het,otal, was selected for use in the following experiments: DZ5'4' 0.78540, which corresponds to 99.9% ethyl alcohol according t o the U. S. Bureai: of Standards table^.^ The normal propyl alcohol was obtained by a similar process from a sarngle of "refined" commercial product,, initial boiling point 96'-98' C: After dist,illation with lime, witth metallic calcium, and fractionation, a liter of the original material yielded 5 7 5 cc with a density of 0.8004! corresponding t o about 993/4y0 alcohol on the basis of Brunel's4 value of 0.7998 25'/4' for 100% and t'he variation per 1% of water of 0.0026 obtained from LandoltBornstein-Roth5. Determination of heat of nziying. The alcohols thus prepared were used in making a series of seven mixtures, which varied systematically in steps 'of approximately 1 2 . 5 7 ~of each component. Thus the first mixture contained by weight 87.507~of ethyl alcohol and 1 2 . 5 0 7 ~of the n-propyl alcohol; the second mixture, 74.93% ethyl and ~ 5 ~ . 0 7 propyl, % etc. About I O O cc of each mixture was made up. I n the course of the preparation of these solutions the heat of mixing was determined a t 25'C. This was accomplished by placing the proper weight of the component which was to be present in greater quantity in any particular case in a half-pint Dewar jar, equipped wit8ha calibrated Beckmann thermometer, a stirrer of the propeller type and a tightly-fitting cork cover t o exclude Bose: Z. physik. Chem., 58, 611 (1907). Gibson, Park8 and Lntimer: J. Am. Chem. SOC.42, 1542 (1920). IT. S. Bureau of Standards. Scientific Paper, No. 197 (1913). Rrunel, Crenshaw and Tobin: J. Am. Chem. SOP.43, 574 (1921). Lrtndolt-Bornstein-Roth: Tabellen, p. 307 (1912).
GEORGE S . PARKS AND JULIUS RAE SCHWENCK
722
moisture. The required weight of the other component was then placed in a half-pint Dewar bottle also equipped with a calibrated Beckmann thermometer, the zero of which was checked daily with that of the Beckmann in the jar. When the two liquids had reached approximately the same temperature, that in the bottle was quickly poured via a thistle-tube lead into the Dewar jar which served as the calorimeter, the mixture was stirred for two minutes a t a rate of 80 R.P.M. and then allowed to stand for a similar period after which the final temperature was taken. By a t once repeating the stirring and waiting interval and again reading the Beckmann, the temperature correction due t o stirring, heat exchanges with the surroundings, etc. was obtained, The corrected temperature change was then used for calculating the heat effect in producing one gram of the solution under consideration. This calculation was made in the usual way, using 6.0 cal. as the heat capacity of the calorimeter and ,582 and .570 cal. as the specific heats at z 5 O C . of pure ethyl and pure 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 the work of Gibson, Parks and Latimer had shown to be justified. The results are given in Table 11.
TABLE I1 Heat of Formation of the Mixtures a t Liquid
$’& ClHbOH
by weight 100.00
87 .50 74.93 62.40 50.02
37.49 25
.oo
12.54 0.00
Mol. fraction of C2H60H I .ooo
25’
Heat, of mixing in calories per mol of per gram of mixture
mixture
. go2
--.036
-I
.796 .684 . 5 66 .439 .303 .158
- ,058 - .079 - .ogz - ,089 - .078 - ,054 ___
___
.000
.’I -2.8 -4.0 -4.8 -4.8 -4.3 -3. I
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. 4” for instance, we shall be referring to the solution containing 62.40% (by weight) of ethyl alcohol. The process of forming the various solutions took place with the absorption of heat and the values in calories for the production of I gm. and I mol of mixture are given in columns 4 and 5 respectively. It will be noticed that our value for liquid No. 5 is - ,092 csl. per gram of mixture, while Bow obtained - , 1 2 5for an almost identical concentration. Certainly it is evident that the hpat effect is very small throughout the entire range of concentrations-an indication that wc are dealing with solutions that are not far from perfect.
PROPERTIES O F MIXTURES O F ALCOHOL
723
Densities and Refractive Indices. The densities of tlie liquids thus made up were next determined in the usual manner, using a specific gravity bottle of about 50 cc capacity. Corrections were made for the buoyancy of the air, and the final values appear in the second column of Table 111; they are probably accurate to the extent of f 0001 gm. per cc. For purposes of comparison determinations were also run on samples of the pure alcohols after they had been put thru the stirring process, etc. in the Dewar jar described in the previous section. This was done because it, seemed inevitable that the samples would absorb traces of moisture during such a calorimetric procedure and we desired to have all our liquids in as comparable a condition as possible. The refractive indices for sodium light were then determined with a ZeissPulfrich refractometer, the method of Moorel for temperature measurement being employed, This ingenious method obviated all necessity of using thermostatic devices and permitted us t o express our results for the standard temperature of 25OC.
TABLE 111 (Temperature, 2 5 O C ) Liquid
Observed
0.7863 0,7886 0.7904 0.7923 0,7942 0.7961 0.7979
Density
Rcfractive Index Observed Calculated
Calculated
--
.3590 .3619 I .3620 3 1.3649 1.3651 I ,3681 I ,3681 4 I .3712 I ,3712 5 6 1.3742 I .3742 1.3772 1,3772 7 8 0.8000 1.3803 I . 3803 0.8018 9 1 ,3833 I n the formation of a perfect solution the resulting volume should be equal t o the sum of the original volumes of the components involved, or in terms of densities the relationship is I
2
I
0.7882 0.7901 0.7921 0.7940 0.7959 0.7979 0 7998
I
’
IO0
I
- =-P,
+ -Pz I
D dl d, where dl and d z are the densities of the components in the pure state, PI and Pzare their corresponding weight percentages in the resulting solution and D is the density of the solution. Using this equation, we obtained the calculated densities appearing in column 3 of the preceding table. These values average less than .03% lower than the observed densities of column 2 , and this fact indicates that the original assumption of “idealness” for t,hese alcohol mixtures is very close to the truth. The refractive indices of the pure alcohols differ by only .0243 and the densities by .o155; hence it was not considered worth while t o apply the 1
Moore: J. Phys. Chem. 25, 281 (1921).
724
GEORGE S. PARKS AND JULIUS RAE SCHWENCK
relationship of Loreirtz and Lorenz' t o this set of liquids. Instead the refractive indices appearing in the fifth column were calcnlated by the siniple equat,ion 100n=Plnl+I-'zn2 where nl and n2 are the indices of refraction for Dhe pure liquids, PI and Pz being their respective weight percentages in any given solution. Comparison of these calculated values with the observed shows that for this system the procedure is fully justified. The straight line relationship between the index of refraction and the weight composition of the solution, coupled with the fact that only an extremely small quantity of liquid is required for the measurement, renders this an excellent Rnalyt,icalmethod for determining the composition of an unknown mixture of the alcohols. The instrument used could be easily read with an accuracy of f I minute, while the two pure alcohols differed by 4'22'; hence analysis to better than . 4 x 8is feasible. Viscosities. Next we ran the viscosities for all the liquids, using an Ostwald viscosimeter in a 2 5' C. thermost,at, regulated t,o ,005'. Thc value of 0.00893 dynes per sq. em., as obtained by Hosking2 was assumed for the water which was used in standardizing this instrument. The time, measured by a stopwatch, averaged around IOO seconds and thus limited the accuracy of these results t o f .00003 dynes per sq. cni.
TABLE IV Viscosities a t 25' C (in dynes per sq. cm.) Liquid
Obscrved values
Calculated (Kendall's equation)
Calculated (Logarithmic equation)
.010go
.or169 .OI233 .o1319 .o I 408 .OI522 . o I 640 .01759 .OI897
.OII57
.OIISI
,01231
.OI22I
~01.314 . o I A04
.012gg .01387 .01488
.01506 ,01621 .01750
. o I 604 .OI738
-_-
The viscosities as determined range from . o ~ o g ofor ethyl alcohol t o .018g7 for the propyl and, while there seems to be no really satisfactory equation for the viscosities of mixtures, the system under consideration provides a means of testing out two of the more promising relationships. Accordingly, in column 3 of Table I V appear the values calculated by use of the cuberoot equation suggested by KendalF : 17%
=x1rIy
+xzqF
Nernst: Theoret,ical Chemist,rg, p. 114(1923). Hosking: Proc. Roy. Soc. N. S.Wales, 43, 37 (1909). a Kendall: J. Am. Chem. SOC.42, 1776 (1920).
42 5
PROPERTIES O F MIXTURES O F ALCOHOL
where ql and q2 are the viscosities of the pure components and XI and X2 are their respective mol fractions. The logarithmic relation, log q =x1 log r]l+XZ log qe was also tried and the results obtained appear in the last column. It i s evident from the data that Kendall's equation is better throughout the entire range of mixtures, being on the average about . 65y0 low while the latter is I .6% low. Boiling Point and Distillate Composition. The boiling points of the various liquids were determined t o the nearest 0 . IO. This wa9 accomplished by placing a 2 5 cc sample in a small distilling flask fitted with a condenser. The liquid was gradually brought t o boiling and, a t the time of delivery of the first drop of distillate, the temperature was read by a calibrated mercury thermometer, graduated t o 0 . I O C . When about I cc of distillate had passed over, the process was stopped and the composition of this distillate determined by measurement of its refractive index. The experimental results appear in columns 2 , 3 and 4 of the following table.
TABLE V Boiling points and Distillate Composition. Barometer, 761 mm Liquid
Boiling Point
78.4OC 79.8 81.4
83.1 85 . o
87. I 89.3 92.7 97.2
CzH60H in distillate Mol fraction
yo by weight 100.0
93 . o 86.0 77.4 68.8 58.3 43.5 22.3 0.0
Part,id Pressure of CzHbOH Observed Calculated
1.000
946 .886 .813 '738 .641 499 ,269
'
'
,000
I n the fifth column we have the partial pressures of ethyl alcohol in the various solutions computed from the observed total pressure (761 nim) and the niol fraction of ethyl alcohol found in the distillate; whilc in the next column are inserted for purposes of comparison the values calculated by the use of Raoult's law. For obtaining these latter results it was necessary to have the actual vapor pressure of pure ethyl alcohol a t each of t,he observed temperatures. This was found by plotting on large coordinate paper the pressures and temperatures as given by the Landolt-R6rnstein-Roth Tabellen (page 387) and then reading off from such a graph the pressure corresponding to the observed boiling point of each solution. Total Pressures and Partial Pressures at 2.5' C . The composition of the vapor phase in equilibrium with the solutions a t 2 5 O C . was determined. This was accomplished by passing air (freed from water and carbon dioxide) through a pair of bubblers, each containing about 2 0 cc of the mixture under
GEORGE S. P A R K S AND JULIUS RAE SCHWENCK
726
consideration. The bubblers were immersed in water baths kept within I O of 2 5 ' C ; closer temperature control is not necessary as the percentage composition of the vapor phase changes very little between 25' and the boiling point. The air thus saturated with the vapor of a mixture was then passed thru a
TABLE VI Ethyl Alcohol in the Vapor of the Liquids a t 25' C. Liquid % 'by weignt MOI fraction of C ~ H ~ O H I
v iAl
A 4
I 0
FIG. I
100.0
I.ooo
92.6
.942 .go2
87.5 79.3 71.5 60.3
a833 .767 .662
45.5 27.0
.3 2 6
0 .o
.ooo
.52I
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 the refractometer method previously described. The results appear in Table VI. Our intention had been to determine also the total pressure of each liquid in the series but, this measurement was postponed t o the very last. Press of other work delayed us for about a year and then, as it seemed possible that the original mixtures niight have undergone some change or contamination during the interval, a new set of four solutions was made up specially for this phase of the work from the pure alcohols1. The difference between the total pressure of the ethyl alcohol and each other liquid in this new series was determined by a differential method, utilizing a glass apparatus similar in form t o the accompanying diagram. Ethyl alcohol, as the reference substance, was placed in the side-arm AI which was closed a t the top by a tested stopcock S1 while the other liquid was placed in the arm A2, being poured in via a similar stopcock Sz. The lower end of the apparatus a t C was connected by pressure tubing with a mercury reservoir which could be raised or lowered as desired and which was carefully manipulat.ed so that a t all times the level of mercury was well abovc
1 For the preparation of these new solutions and the subsequent measurement of their respective vapor pressures, we are indebted to Mr. 0. R. Lovekin of this laboratory.
PROPERTIES OF MIXTURES OF ALCOHOL
727
the fork B and thus the two liquids were kept entirely separate. The procedure was to introduce by the two stopcocks 5 cc. samples into each arm of the apparatus, then t o lower the mercury reservoir until a small vapor phase was produced and wait a few minutes. During this time much of the air dissolved in the liquids was evolved and could be forced out of the apparatus by raising the mercury level and simiilt,aneously opening the two stopcocks t o permit of its escape. By repeating this operation two or three times all the air originally in the aamp!es could easily be removed. Then the mercury reservoir was so adjusted as t o permit the existence of a small vapor phase and the difference in mercury levels in the two tubes was measured by a cathctometer a t quarter-hour intervals, the temperature registered by a thermometer placed between the two arms being noted a t the time of each reading. Corrections, when necessary, were made for any differences in the columns of alcohol liquid in the two arms. These measurements were all made in a small room in which the temperature could be kept very close t o 2 5 O C and any departure from this standard point (seldom more than a few tenths of a degree) could be easily corrected for by a rough determination of the rate of change of this pressure difference with temperature in each particular case. Regnault’s value of 59.0 mm. as corrected by Bunsen1 was taken as the vapor pressure of the ethyl alcohol and, by means of the differences measurcd in the manner just described, the vapor pressure of each of the five remaining liquids was computed; the results appear in the following table, column 4.
TABLE VI1 Data for New Set of Mixtures Liquid
Mol fraction of (IzHsOH in liquid
Mol fraction of C~HSOH in vapor phase .ooo
I
1.000
A
.836 .656
-924 .823
459
B C
I
D
.241
.678 ,440
9
.000
,000
*
Vapor pressure of liquids Experimental “Ideal”
59. o (assumed) 53.omm 53.1 mm 46.7 ’) 47.5 ” 40.4 )’ 39.7 9 , 32.3 31.8 23.2
The values (column 3) for the composition of the vapors in equilibrium with these new solutions were not determined directly but were obtained graphically from a smooth curve representing the data of Table VI. In this sense they are really experimental and, when multiplied by the correfponding total pressures for the various solutions, give us the observed partial pressures of ethyl alcohol appearing in column z of Table VIII. The observed partial Landolt-Rornstein-Roth: Taballen, p. 386 (1912).
728
GEORGE S. PARKS AND JULIUS RAE SCHWENCK
pressures of normal propyl alcohol were obtained in similar fashion. The "ideal" pressures in all cases were derived on the assumption of Raoult’s law: PA = NAPA where PA and NA are respectively the parbial pressure and the mol fraction of component A in a given solution and PAis itjsvapor pressure in the pure state.
TABLE VI11 Partial Pressures of the Coniponents at 25’C. Tiquid
Partial preasure of C2ITaOH Observed Ideal
mm
---
4.0
’,
3 . 8 mm
8.4
”
”
13.0
”
”
18.0
’,
23.2
”
I
59.ornm
--
A
49 .o
B
39.1
,,
49.3mm 38.7 ,,
C D
27.4 14.3
”
27.1
71
14.2
9
0.0
”
,,
Partial pressme of C3P~OI-l Observed Ideal 0.0
8.0
12.6 17.6
,, ” ”
--
It will be noticed that on the whole the values derived from our experimental data are somewhat higher than those called for on the assumption that we are here dealing with ideal or perfect solutions. This deviation, however, is the usual one when heat is absorbed on the formation of a mixture from the pure components. The Entropy of an Egui-Molal Mixture. In this connection it is interesting to check u p on an assumption made by Gibson, Parks and Latimer. They needed the entropy change for the process I/Z mol EtOH+ 1 / 2 mol PrOH = I niol of mixture and, in the absence of the proper data, obtained the value I .37 cal. per degree by assuming that the two alcohols form a perfect solution. We now are in a position t o calculatc the entropy change by means of the thermodynamic equation AF=AH-TAS . By plotting our figures for the partial pressures of the two alcohols against the corresponding mol fractions, drawing smooth curves and reading from these the partial pressures of each in an equi-molal mixture, we obtain 2 9 . 8 mm and 12.2 mm as the respective values for ethyl and n-propyl alcohols. From these AF, the free energy change a t 2 j°C for the above procms, can be found : 29.8 12.2 AFZ50C= 298R,[$ln-+$ln] = -393 cal. 59.0
23.2
Then from our data on the heat of mixing, we obtain 5 ml. as the increase in heat content for the production of the mol of niixturc. Hence, ASZ5oc=
393+5= I .34 cal. per degree, a result 298
on the assumption of a n ideal solution.
alniost identical with that calklated
PROPERTIES O F MIXTURES O F ALCOHOL
729
Summary Reviewing the results of the various measurements, we find that (I) -4very small heat absorption (in no case more than 5 calories per mol of resulting mixture) takes place on formation of the several solutions. ( 2 ) An extremely small volume shrinkage-on the average only .025y0 -accompanies the process. ( 3 ) The refractive indices of the various liquids are practically a straight line function of their weight compositions. (4) The observed viscosities foy the various solutions exhibit on the average a positive deviation of 2 / 3 y 0 from Kendall’s “cube root” equation. ( 5 ) The partial pressures of each component as calculated from the experimental data a t 25’ C are in general slightly higher than those called for on the basis of Raoult’s law. Judging these facts as a whole, we may conclude that the system under consideration is almost “perfect”. Department of Chemistry Stanford Universi:y, California Mnrch 10, 1924.
