EQUILIBRIUM DISTRIBUTION O F ACETIC ACID BETWEEN ISOPROPYL ETHER AND WATER Depaitment of
ANDREW A. SMITH AND JOSEPH C . ELGIN Chemical Engineering, Princeton University, Princeton, New Jersey Received December 18, 1934
The removal of undesirable constituents or the recovery of valuable components from solutlbn by selective extraction with non-miscible solvents is a potentially important chemical engineering operation for which few significant quantitative data are available. I n the course of experimental studies of the rate of extraction of acetic acid from aqueous solution by isopropyl ether, comprising one phase of an extensive program of investigation of the rates of liquid-liquid extraction processes in progress in the Department of Chemical Engineering of Princeton University, it became necessary to have available distribution data for this system. Inasmuch as the literature apparently reveals no published measurements of this distribution equilibrium, our data on it are presented in this paper. As this system is the basis of a commercial process for the recovery of acetic acid from aqueous solution by extraction, these results are of industrial as well as academic interest. EXPERIMENTAL TECHNIQUE
Material. The measurements were made with purified isopropyl ether from the Eastman Kodak Company and with distilled water as the solvents, and also with technical isopropyl ether obtained directly from the Carbide and Carbon Chemicals Corporation and with the local tap water. The technical ether was dry to anhydrous copper sulfate, contained an acidity less than 0.01 per cent, and less than 3 per cent by weight of isopropyl alcohol. The acetic acid was glacial (99.5 per cent) acid of the U.S.P. grade. Analytical method. Acetic acid concentrations in both water and ether phases were determined volumetrically by titration of IO-cc. to 25-cc. samples with a standard alcohol solution of sodium hydroxide. This was prepared by dissolving the usual purified stick caustic in 95 per cent ethyl alcohol, standardized against constant-boiling hydrochloric acid, and employed in concentrations of 0.1 N , 0.5 N , and 1.0 N according to the concentration range in the system under investigation. Prior to titration the samples were diluted to three times the original volume to insure com1149
1150
ANDREW A. SMITH AND JOSEPH C. ELCTIN
plete miscibility with the titrating solution, alcohol acting as a common solvent. The indicator was phenolphthalein. Procedure. The distribution measurements were carried out by the ordinary laboratory method usually employed for this purpose. The two solvents, water and isopropyl ether, one containing a sufficient quantity of acetic acid to produce the final acid concentration range desired, were placed in 250-cc. glass-stoppered bottles. The volume of each solvent was usually 100 cc. In order to approach the equilibrium from both sides and thus minimize possible error due to failure to attain equilibrium, independent series of experiments were conducted with acetic acid initially in the water phase and in the ether phase, respectively. After sealing, the bottles were vigorously agitated, placed in a thermostat at 20 f 0.5"C., and allowed to remain for a period of 170 hours wIth vigorous agitation at frequent intervals. Ten-cc. samples from each layer were analyzed for acid after 48 hours and 25-cc. samples at the end of 170 hours. Inasmuch as the analyses of the two sets of samples checked closely, it wasassumed that equilibrium had been attained. Precaution was taken to prevent evaporation of ether from the samples, these being run directly into alcohol in stoppered flasks. As a check, the sum of the quantities of acid found in each layer by the analyses was compared with the known total quantity of acid originally added to the system. Duplicate experiments were performed for the majority of concentrations. EXPERIMENTAL RESULTS
The results of the measurements for various acid concentrations are listed in table 1. In this table equilibrium concentrations, expressed in gram-moles per liter of solution, for the water, C,, and ether, C,, phases are given in columns 1 and 2, respectively. Corresponding values of the apparent distribution coefficient, Dobsd. = C,/C,, calculated from the observed concentration on the assumption that the simple ideal distribution law for single undissociated molecules is obeyed, are given in column 3. It is evident from the data of table 1 that the ideal distribution law is only approximately obeyed, inasmuch as the apparent distribution coefficient calculated from the ratio CJC, slowly increases with total acid concentration. This is an expected result, since one is undoubtedly not dealing with a single molecular species in both phases, acetic acid probably undergoing increasing dissociation in the aqueous phase with increasing dilution and association in the ether phase. Especially at the higher concentrations the apparent values of the coefficient where the acid was initially in the ether phase are consistently somewhat higher than for corresponding concentrations where the acid was present initially in the water phase. This apparently indicates that true equilibrium was not entirely attained. However, analyses of both phases in each experiment over a 170-hour period with frequent agitation gave no
1151
EQUILIBRIUM DISTRIBUTION OF ACETIC ACID
further detectable change in observed concentrations; consequently, further possible change toward equilibrium would be infinitely slow. The apparent variation from true equilibrium is less significant than at first appears, and is probably within the limits of possible analytical error. A deviation in the observed concentrations from the true equilibrium values is multiplied twofold in the calculation of the distribution ratio, and, furthermore, an average error of 0.02 to 0.05 cc. in the titration value could account for the minor discrepancies obtained. TABLE 1 Distri'bution of acetic acid between water and isopropyl ether at 80°C. ~
EQUILIBRIUM CONCENTRATION I N ORAM-POLE8 PER LITER
Ether, C,
I
Water, Cw
~
DIS'TRIBUTION COEFFICIQNT,
Observed
1
~~
D
=
c6/c,.,
Calculated from equation 3
(a) Pure ether and distilled water
0.0136 0.0140* 0.0254 0,0560* 0.0577 o.ii86 0.1430* 0.1733 0.2290* 0.2803 0,3260* 0.3994 0.6740
0.0732 0.0785 0.1450 0,2980 0.3030 0.6025 0.6925 0.8900 1.1070 1.3600 1.4300 1 ,7970 2.7800
0.185 0.178 0.177 0.188 0.190 0.196 0.206 0.196 0.207 0.206 0.228 0.222 0,242
0.180 0.180 0.182 0.186 0.187 0.194 0.197 0.202 0.208 0.215 0.217 0.226 0.253
(b) Technical ether and tap water
0.0543 0.1480 0.2740
0.2960 0.7720 1 ,3450
0.200 0.192 0,232
* I n these experiments acid was added initially to the ether.
The results with tap water and technical isopropyl ether are slightly higher than those for the pure solvents. This is probably to be accounted for mainly on the basis of the isopropyl alcohol present in the technical ether. In figure 1 is plotted the actual distribution equilibrium curve for the range investigated. Distribution data have usually been represented by an empirical equation of the form The present equilibrium curve may be approximately represented over the concentration range investigated by such a parabolic equation, since a
1152
ANDREW A. SMITH AND JOSEPH C. ELGIN
satisfactory straight line is obtained in a log-log plot of C , against Cw. The slope of this line is a = 1.09, and its intercept, D’ = 0.148. Therefore C,/C~oe= 0.148 (1) which equation may be employed for the approximate calculation of the distribution of acetic acid between water and isopropyl ether over the present concentration range. DISCUSSION
It has been pointed out by Almquist (1) that the apparent distribution coefficient may frequently be expressed most accurately as a function of
CONCN N ISOPROPYL E W E R
FIG. 1
I1PtES/IIlCR
c.
rlOLES/LIlCR
FIG.2 CURVEFOR THW DISTRIBUTION OF ACETICACID BETWEEN ISOPROPYL ETHERAND WATER
FIQ. 1.
EQUILIBRIUM
FIU. 2.
VARIATION OF THE
APPARENT DISTRIBUTION COEFFICIENT
WITH CONCEN-
TRATION I N THE AQUEOUS PHASE
0, acid added initially t o the water; @, acid added initially to the isopropyl ether.
the concentration in the water phase by a simple linear equation of the form Depparent = C e / C w = Ds C w Di (2)
+
where D, is the slope and Di the intercept. Such an equation possesses theoretical significance and supplies an accurate and useful means of recording, testing, and interpolating distribution data. That the data for the present case conform to such an equation may be seen from figure 2, in which a straight line results from a plot of the ratio CJC, against Cw. The slope of this line is 0.027 and its intercept is 0.178. Hence, its equation is
DsPParent = CJC, = 0.027Cw + 0.178
(3)
This equation may be empIoyed with considerable accuracy for distribution calculations for acetic acid between water and isopropyl ether and is
EQUILIBRIUM DISTRIBUTION OF ACETIC ACID
1153
recommended for this purpose. In column 4 of table 1 values of the apparent coefficient calculated from the concentrations in the water phase by means of equation 3 are given. Since equation 3 is based on the measurements approaching the equilibrium from both sides, thus representing average values, it is probable that the calculated values in column 4 represent the true distribution ratio more accurately than do the individually observed values in column 3. Based on theoretical considerations involving the possible distribution and association equilibria existing in such a case and assuming no association in the water phase, Almquist (1) has derived, as a general theoretical relation for the distribution of an organic solute between water and an organic solvent, the equation:
+
C,/Cw = nK&;C:-' K2 (4) In this equation n = number of molecules in the polymer in the non-aqueous phase. K 3 = (A:)/(A')" = association constant in the non-aqueous phase, i.e., the equilibrium constant for the reaction , nA' = A:. K z = A'/A = distribution constant for single molecules between the two phases. A' and A = concentrations of single molecules in the non-aqueous phase x and the water phase w, respectively. If the solute is mainly associated only into double molecules in the now aqueous phase, then n = 2, equation 4 reduces to and the apparent distribution coefficient is a linear function of C,. Since this is the case for acetic acid between the present solvents, it is evident that acetic acid is associated into double molecules in isopropyl ether over the concentration range investigated. In the present case the slope of equation 3, 0.027, is identical with 2K&;, the slope of equation 5, and the intercept, 0.178, corresponds to Kz. It therefore follows that the value of the distribution constant for single molecules of acetic acid between water and isopropyl ether is Kz = 0.178, and the value of the association constant for acetic acid in isopropyl ether is Ks = 0.43, concentrations being expressed in moles per liter. At very low concentrations, owing to dissociation in the water phase, and at very high concentrations where association in the water phase may become appreciable, simple relations of the above form may be expected to be no longer valid. REFERENCE (1) ALMOUIST:
J. Phys. Chem. ST, 991 (1933).
