Ind. Eng. Chem. Res. 1994,33, 1373-1379
1373
Water-Enhanced Solvation of Organic Solutes in Ketone and Ester Solvents Jane H. Lee,+Vincent Van Brunt,$ and C. Judson King' Department of Chemical Engineering and Lawrence Berkeley Laboratory, University of California, Berkeley, California 94720
Previous research has shown that the solubilities of dicarboxylic acids in certain electron-donor solvents are substantially increased in the presence of water. Information on solubilities, liquidliquid equilibria and maximum-boiling ternary azeotropes was screened so as to identify other systems where codissolved water appears to enhance solvation of organic solutes in solvents. Several carboxylic acids, an alcohol, diols, and phenols were selected for examination as solutes in ketone and ester solvents. Effects of water upon solute solubilities and volatilities were measured. Results showed that water-enhanced solvation is greatest for carboxylic acids. Solute activity coefficients decreased by factors of 2-3,6-8, and 7-10 due to the presence of water for mono-, di- and tricarboxylic acids, respectively. Activity coefficients decreased by a factor of about 1.5 for ethanol and 1,2propanediol as solutes. Water-enhanced solvation of phenols is small, when existent.
Introduction Solubilities of certain organic solutes in some organic solvents are much increased in the presence of codissolved water. The solubility of the solute in the water-containing solvent can even substantially exceed the solubility in water alone. Such is the case for fumaric, succinic, and adipic acids in ketones and several other electron-donor solvents (Starr and King, 1992a). For example, the solubility of adipic acid in water-saturated methylcyclohexanone at 25 "C is 9.7% (w/w), far greater than the solubility in either anhydrous methylcyclohexanone (1.6 % ) or water itself (1.4%) (Starr and King, 1992a). Water-enhanced solvation is a manifestation of a phenomenon where the intermolecular interactions in a solute-solvent system containing water are stronger than in the water-free solute-solvent binary system. The implication is that water and the solute interact more strongly than do the solute and solvent. In cases such as the adipic acid-water-methylcyclohexanone system, where the solubility of a solute is greater in a water-containing solvent than in water or solvent alone, the further implication is that the solute, water, and solvent form a ternary solvate or complex in which the interactions are stronger than in the solute-water and solute-solvent binaries. This conclusion can be generalized to any case where the activity coefficient of a solute reaches a lower value in a solvent-water mixture than in water or solvent alone. On the basis of related spectroscopic evidence it has been proposed (Starr and King, 1992a) that the complex or solvate formed in systems of dicarboxylic acids, water, and electron-donor solvents takes the form of water molecules positioned between the acid and the solvent. The oxygen atoms of water would then hydrogen bond with the available protons of the carboxylic acid, and the hydrogen atoms of water would form hydrogen bonds with molecules of the electron-donor solvent. If water serves to decrease the activity coefficient of the solute in the solvent sufficiently, that effect can be used as a basis for a separation process. Starr and King (1992a,b) proposed a processing approach wherein a
* To whom correspondence should be addressed.
Present address: General Mills, Inc., West Chicago, IL 60185. Present address: Department of Chemical Engineering, University of South Carolina, Columbia, SC 29208. t
solvent, such as a cyclic ketone, is used for extraction of a solute, such as one of the three aforementioned dicarboxylic acids. Water is then removed selectively from the extract, e.g., by stripping, thereby causing precipitation of the solute and regeneration of the solvent for reuse. The solvent resaturates with water upon return to the extractor and thereby regains solvent capacity. The advantage of such a regeneration process lies in the removal of a minor component (water) from the extract, rather than evaporation of the entire solvent, to accomplish the precipitation and regeneration. Other types of separation could be based upon changing the water content of a solute-water-solvent solution selectively, if doing so substantially alters the activity coefficient of the solute. One example would be increasing the capacity of a solvent for leaching solutes from a sorbent by adding water to the solvent. Another would be enhancing adsorption or membrane permeation of a solute from a solvent by selectively removing water from the feed solution. Separations of this sort are not limited to solutes of low solubility. The goals of the present work were to use several screening criteria to identify other classes of organic solutesolvent systems for which water-enhanced solvation appears to occur and to measure the sizes of the waterenhancement effects in representative systems from among the classes identified.
Screening of Water-Containing Ternary Systems The screening criteria were selected to reveal watercontaining systems having preferential ternary interactions that produce a lower activity coefficient of the solute in a homogeneous mixture of water and solvent than in water or solvent alone. Systems meeting three different criteria were identified (1) Ternary systems in which a solute exhibits a maximum solubility at an intermediate composition of solvent and water. (2) Maximum-boiling ternary azeotropes. (3) Ternary liquid-liquid systems in which a relatively small amount of the third component serves to create miscibility in an otherwise highly immiscible binary system. All three of these phenomena reflect preferential, positive interactions in the ternary system. Table 1 lists ternary solute-water-solvent systems for which a maximum solubility of the solute in an interme-
0888-5S85/94/2633-1373$04.50~0 0 1994 American Chemical Society
1374 Ind. Eng. Chem. Res., Vol. 33, No. 5, 1994 Table 1. Solute-Water-Solvent Systems for Which a Maximum Solute Solubility Exists in a Mixture of Water and Solvent solute solvent SmUJSatlh reference A. Systems Where Water and the Solvent Show Limited Miscibility i. Solid-Liquid-Liquid Triple Points Exist cyclohexanone 8.6 Starr and King, 1992a fumaric acid methylcyclohexanone 7.5 Starr and King, 1992a methyl isobutyl ketone 5.7 Starr and King, 1992a 2-heptanone 5.2 Starr and King, 1992a 3-heptanone 3.8 Starr and King, 1992a 4-heptanone 4.1 Starr and King, 1992a acetophenone 5.5 Starr and King, 1992a n-butyl acetate 5.7 Starr and King, 1992a di-n-butyl ether 1.6 Starr and King, 1992a Starr and King, 1992a 1-octanol 1.6 tritolyl phosphate 1.7 Starr and King, 1992a cyclohexanone 7.7 succinic acid Starr and King, 1992a methylcyclohexanone 6.3 Starr and King, 1992a methyl isobutyl ketone 5.0 Starr and King, 1992a 2-heptanone 4.1 Starr and King, 1992a 3-heptanone 2.7 Starr and King, 1992a 4-heptanone 3.2 Starr and King, 1992a acetophenone 4.6 Starr and King, 1992a di-n-butyl ether 3.6 Forbes and Coolidge, 1991 1-butanol 3.2 Hafez and Hartland, 1976 adipic acid cyclohexanone 6.5 Starr and King, 1992a methylcyclohexanone 6.0 Starr and King, 1992a methyl isobutyl ketone 3.4 Starr and King, 1992a 2-heptanone 4.0 Starr and King, 1992a 3-heptanone 3.2 Starr and King, 1992a 4-heptanone 3.6 Starr and King, 1992a acetophenone 5.1 Starr and King, 1992a thiodiacetic acid (from aqueous salt solution) 2-butanone 2.8 Clegg and Bearse, 1950 succinicacid m-nitrophenol pnitrophenol
ii. No Triple Points Exist cyclopentanone
6.6
B. Systems for Which There Is an Internal Binodal Curve ethanol 1.05 ethanol 1.2
C. Systems for Which There Is No Region of Liquid Immiscibility acetone 5.4 1-propanol 3.1 ethanol 1.7 methanol 1.1 dioxane 1.3 adipic acid acetone 4.5 ethanol 1.5 picric acid ethanol 1.7 1-propanol 3.3 2-propanol 1.5 5-methylhydantoin ethanol 5.3 5-ethylhydantoin ethanol 2.8 5-isobutylhydantoin ethanol 1.6 2-methanamidobutanoic acid 2.0 ethanol 1.2 2-methanamido-4-methylvalericacid ethanol 1.2 acetanilide ethanol succinic acid
diate composition of solvent and water has been reported. The ratio of the maximum solute solubility (S,) to that in the anhydrous solvent ( S d ) is also given. For systems where the water-solvent binary is not fully miscible and solid-liquid-liquid triple points exist, the ratio reported is that of the solubility of the solvent in water-saturated solvent to the solubility in the anhydrous solvent. The temperature in each case is the closest temperature to 25 "C for which data are reported in the original source indicated. Not included in Table 1are numerous ternary systems showing upward curvature (i.e., bowing toward the solute apex) in the solid-liquid equilibrium relationship on a triangular diagram but not exhibiting internal maxima
Bertini and Pino, 1959 Duff and Bills, 1930 Duff and Bills, 1930 Bancroft and Butler, 1932 Bancroft and Butler, 1932 Bancroft and Butler, 1932 Bancroft and Butler, 1932 Herz and Lorentz, 1929 Bancroft and Butler, 1932 Bancroft and Butler, 1932 Duff and Bills, 1931 Duff and Bills, 1931 Duff and Bills, 1931 McKeekin et al., 1935 McKeekin et al., 1935 McKeekin et al., 1935 McKeekin et al., 1935 McKeekin et al., 1935 Gregg-Wilson and Wright, 1928
(Francis, 1941;Gregg-Wilson and Wright, 1928). Many of these systems are chemically similar to those listed in Table 1. The solutes in the systems identified all have electronacceptor (proton-donor) characteristics and include carboxylic acids, phenols, hydantoins, and an anilide. The greatest solubility enhancements occur for carboxylicacids. The solvents in the systems identified all have electrondonor characteristics and include ketones, esters, ethers, alcohols, and a phosphate. The greatest solubility enhancements due to water occur for esters and ketones, particularly cyclic ketones. Table 2 reports three maximum-boiling ternary azeotropes containing water. These cases involve phenols as
Ind. Eng.Chem.Res., Vol. 33, No. 5,1994 1375 Table 2. Maximum-Boiling Ternary Azeotropes Containing Water reference azeotropic system water/phenol/acetone Zernicke (1995) water/phenol/pyridine Seidell and Linke (1952), Mertzline (1936) water/phenol/ Seidell and Linke (1952), Zhuravlev and Bychkova (1947) m-phenylenediamine Table 3. Water-ContainingSystems Having Unusually Shallow but Wide Binodal Curves component promoting incompletely miscible components4 miscibility phenol formic acid water aniline formic acid water ethyl acetate acetic acid water acetic acid water m-toluidine phenol trichloroacetic acid water lactic acid water aniline methanol water isobutanol methanol water aniline methanol water p-cresol ethyl acetate ethanol water n-butanol ethanol water isobutanol ethanol water water phenol pyrogallol phenol water aniline n-butanol acetone water isobutanol acetone water water acetone glycerol tert-butanol water n-butanol 4 All data taken from Francis (1941)except the last entry, which was taken from Francis (1967).
the electron acceptor or proton donor and a ketone, a pyridine, and an amine as electron donors. Table 3 presents water-containing systems identified as having unusually shallow but wide binodal curves, thereby showing that the third component serves to create specificternary attractive interactions. The criterion used to identify these systems was that the width of the binodal curve, in the absence of the component promoting miscibility, is at least 4 times the height of the binodal curve, in consistent units. These systems are composed of carboxylic acids, alcohols, phenols, and in one case, glycerol as proton donors, and anilines, ketones, esters, alcohols, and phenols as proton acceptors. All systems presented in Tables 1-3 are consistent with the concept postulated by Starr and King (1992a), that the specific ternary interactions come from water being placed between electron-acceptor (proton-donor) solute molecules and electron-donor solvent molecules in a ternary complex or solvate, with the water serving both as an electron donor to hydrogen bond with the solute and as an electron acceptor to hydrogen bond with the solvent. The implication is that the ternary complex or solvate is energetically more stable than the binary complex or solvate between the solute and solvent alone. A similar complex has been inferred from spectroscopic measurements of systems of phenols in water-containing ester and ketone solvents (Smol'skaya et al., 1991; Smol'skaya and Egutkin, 1994). The largest effects seem to occur with ketone and ester solvents. On that basis and on the basis of the insight provided by Tables 1-3, acetic, propionic, citric, and gallic acids, ethanol, 1,2-propanediol,2,3-butanediol,and various phenols were chosen as solutes for further study, with methylcyclohexanone, cyclohexanone, methyl isobutyl ketone, and n-butyl acetate as solvents. These were felt to be systems that could display water-enhanced solubility and could be of interest for large-scale separations.
Experimental Technique Two methods were used to determine the extent of water-enhanced solvation in these systems (Lee and King, 1993). Measurement of the enhancement of solubility was used for solutes of low enough solubilities, and measurement of the reduction of equilibrium partial pressure over a solution in solvent (vapor headspace analysis) was used for sufficiently volatile solutes. Solvents and liquid solutes were contacted with well regenerated Davison Chemical 4A molecular sieves when it was desirable to reduce the water content of the solvent or solute as received. For both solid-liquid and vapor-liquid equilibrium measurements, the concentration of water in the liquid organic phase was determined by Karl Fischer titration, replacing the standard methanol solvent with a modified 2-methoxymethanol GFS Karl Fischer reagent, which prevents formation of the acetals, ketals, and accompanying water that otherwise interfere with analyses in carbonyl-containing solutions. Additional details on the experimental procedures are given elsewhere (Lee and King, 1993). The methylcyclohexanone used was reported by the vendor (Fluka) to be 68 % 3-methylcyclohexanone and 31.4 % 4-methylcyclohexanone, 99.4% purity overall. Solubility Measurements. Solvent-water mixtures were prepared by adding measured volumes of solvent and water to 20-mL scintillation vials. The solid solute of interest was added in excess, and the vial was then placed in a shaker bath for 12 h or more at an elevated temperature, following which the vial was placed in another bath at 25 "C for 24 h to reach equilibrium. The supernatant solution was then analyzed by gas chromatography (GC), high-performance liquid chromatography (HPLC), or two-phase back titration, as appropriate to the solute. GC was used for phenol and xylenols, with a Porapak Q (Waters Associates) column at 225 "C. HPLC was used for higher boiling phenol derivatives and citric acid, with a Cle Resolve column (Waters Associates) and a mobile phase consisting of a 1:l v/v ratio of acetonitrile and 2% (w/w) aqueous acetic acid. Two-phase back-titration for acids was carried out with 0.1 N aqueous NaOH solution. Vapor-Headspace Measurements. Measured amounts of solute, solvent, and water were added to 50mL vials, leaving about 30 mL of vapor space. Teflonlined silicone septa were applied to the vials by pressure from a crimper. The vials were placed in a shaker bath at 25 "C for at least 24 h, with the water level of the bath high enough to cover the liquid and vapor portions of the sample completely. Samples of the vapor were taken with a water-jacketed l-mL gas-tight syringe, with the jacket temperature set 3 "C higher than the sample temperature. During sampling of the vapor phase the syringe needle penetrated the septum, with the tip only a few mm away from the liquid surface within the vial. The liquid phase was sampled with an ordinary 1-pL syringe. Analyses of both the liquid and vapor samples were made by GC with a Porapak PS (Waters Associates) column and a flame-ionization detector. Carboxylic acids form dimers in the vapor phase. Corrections were made for this effect by means of the technique described by Sebastiani and Lacquaniti (1967), using information on dimerization of acetic and propionic acids (Prausnitz and Tsonopoulos, 1970). Series of measurements were carried out at fixed mole ratios of solute to solvent, with varying amounts of water added.
1376 Ind. Eng. Chem. Res., Vol. 33, No. 5,1994 Table 4. Values of Enhancement Factor, E, Determined by Vawr-HeadsDace Techniaue. 25 ' C _ _ _ ~
n o
....
0.6
1
0.4
1
0.2''
+
~
......+
"
0.00
WCHW = 0.02 [mol/mol] AACHM = 0.04 [mollrnol]
"
"
'
"
x
I
'
0.08
0.04
0.12
[mol/mol]
Figure 1. Liquid-phase activity coefficients for acetic acid (AA) in cyclohexanone(CHEX) at 25 "C, as influenced by the mole fraction of water in the liquid phase. 0.02 and 0.04 mol acetic acid/mol cyclohexanone.
.
1.2
solutesolvent y, enhancement solute solvent? mole ratio high/lowb factor, E acetic acid CH 0.02 1.2 0.015 acetic acid CH 0.04 2.8 0.14 propionic acid CH 0.06 2.0 0.25 propionic acid MCH 0.05 2.6 0.19 ethanol MCH 0.02 1.4 0.025 ethanol CH 0.02 1.5 0.024 ethanol CH 0.05 1.5 0.093 1,2-propanediol MCH 0.04 1.6 0.055 a CH = cyclohexanone, MCH = methylcyclohexanone. b Ratio of highest ydlowest yI. 2.0 c
....
0.6
1
0.41 0
"
0.02
"
"
0.04
0.06
'
I
0.08
x w[mol/mol] Figure 2. Liquid-phase activity Coefficients for propionic acid in cyclohexanoneat 25 OC, as influenced by the mole fraction of water in the liquid phase. 0.06 mol propionic acid/mol cyclohexanone. 1.2 I
I I
c
0.3 I 0
"
'
'
' '
"
'
'
'
0.09 x w [rnol/rnol]
'
"
I 0.18
' ' '
Figure 3. Liquid-phase activity coefficients for propionic acid in methylcyclohexanone at 25 OC, as influenced by the mole fraction of water in the liquid phase. 0.05 mol propionic acid/mol methylcyclohexanone.
Experimental Results Carboxylic Acids. Figures 1,2,and 3 report activity coefficientsdetermined by the vapor-headspacetechnique for acetic acid in cyclohexanone, propionic acid in cyclohexanone, and propionic acid in methylcyclohexanone, respectively. Solute activity coefficients are plotted vs. the mole fraction of water in the mixture, a t fixed molar ratios of solute to solvent. Since the solute and water are both relatively dilute, the molar ratio of water to solute is derivable from the ratio of the horizontal coordinate to the fixed molar ratio of solute to solvent. Original data for these and subsequent results are given elsewhere (Lee and King, 1993). As expected, the activity coefficients of acetic and propionic acids in these cyclic ketone solvents decrease substantially with increasing water content, thereby confirming water-enhanced solvation. It is interesting to note in Figure 3 that the activity coefficient levels off above xw = 0.10, i.e., above xwlxs = 2.0. This phenomenon is consistent with the hypothesis put forward by Starr and King (1992a1, based upon earlier work of Christian, Affsprung, and associates (Christian et al., 1965;Wood et al., 1966;Van Duyne et al., 19671,that two water molecules
0.0
0.04
w w Wgl
0.08
0.12
Figure 4. Solubility of citric acid in methylcyclohexanone at 25 OC, as influenced by the weight fraction of water in the liquid phase. (Mole fraction = 6.2 X weight fraction.)
are associated with each carboxylic group in the ternary complexes or solvates that are formed. The data in Figure 2 do not extend to xw/x, = 2.0, so it cannot be determined whether the activity coefficient would also level off beyond that ratio of water to acid for the propionic acidcyclohexanone system. The data in Figure 1 for acetic acid in cyclohexanone at a molar ratio of 0.04 do extend above xw/xs = 2.0, corresponding to xw = 0.08, but the scatter of the data precludes a good test of whether or not yalevels off above that point. Less decrease of yefor xwlxs > 2.0 may explain the much lesser slope of the data for so1ute:solvent mole ratio = 0.02 in Figure 1,since most of the data for that mole ratio are in that range. To compare the relative magnitudes of the waterenhancement effects, it is appropriate to normalize the data so as to take into account differences in the activity coefficients in the absence of water and differences in the solutesolvent mole ratio. Following a linear-free-energy concept and postulating that the mole ratio of water to solute is the dominant independent variable, an enhancement factor, E, for comparing data in different systems and at different mole ratios was defined as -A(log yB)/ (xwlxs)in the region of initial decrease of ye,i.e., at low water content. Values of E were determined by linear regression of log ysvs xwIxB(Lee and King, 1993). Table 4 gives values of E derived for the different systems examined, as well as the ratio of the solute activity coefficients measured at the lowest and highest water contents examined. The values of E are similar for the two propionic acid systems and the acetic acid system at solutesolvent ratio = 0.04. Figure 4 displays the measured solubility of citric acid in methylcyclohexanone a t 25 "C,as a function of the weight fraction of codissolved water, extending up to the triple point at which there is equilibrium with both an aqueous phase and a solid citric acid phase. Table 5 lists measured solubilities of citric acid in methylcyclohexanone and n-butyl acetate and of gallic acid (3,4,5-trihydroxybenzoic acid) in cyclohexanone and methylcyclohexanone
Ind. Eng. Chem. Res., Vol. 33, No. 5, 1994 1377 Table 5. Enhancements of Solubilities of Carboxylic Acids due to the Presence of Water, 25 O C w t fraction
solute solventa citricacid MCH citric acid
BA
gallicacid
CH
gallic acid
MCH
water 0.002 0.11 0.001 0.013 0.01 0.13 0.006 0.079
solubility of solubility solute (wt fraction) ratiob 0.25 1.70 6.9 0.0090 0.087 9.7 0.69 0.75 1.1 0.52 0.83 1.6
a BA = n-butyl acetate, CH = cyclohexanone, MCH = methylcyclohexanone. Ratio of solute solubility at higher water content to solubility at lower water content.
*
a
+ *
ET/CHEX = 0.05 ET/CHEX = 0.02 ET/MCHEX = 0.02
3.4 Y
0
0.04
0.08
0.12
xw [rnol/mol]
Figure 5. Liquid-phase activity coefficients for ethanol (ET) in cyclohexanone (CHEX) and methylcyclohexanone (MCHEX) at 25 "C, as influenced by the mole fraction of water in the liquid phase. 0.02 and 0.05 mol ethanol/mol cyclohexanone;0.02 mol ethanol/mol methylcyclohexanone.
at 25 OC and at both low and high water contents. In all cases in Table 5 except for citric acid in n-butyl acetate, the higher water content corresponds to the triple point, i.e., saturation with water. It can be seen that the enhancement of the solubility of citric acid due to water is very large, as could be anticipated from the fact that it is a tricarboxylic acid. Indeed, putting these results together with those of Starr and King (1992a) for dicarboxylic acids, the solvationenhancement effects in ketone and ester solvents are in rough proportion to the number of carboxylic groups per molecule-by factors of 2-3 for monocarboxylic acids (acetic and propionic), 6-8 for dicarboxylic acids (adipic,fumaric, and succinic), and 7-10 for the tricarboxylic acid (citric). The increases in solubility of gallic acid are much less. The amount of water present at saturation is only 0.9-1.6 mol/mol gallic acid in the solvent phase. There does not appear to be extra water enhancement of solvation associated with the three hydroxy groups. Ethanol. Figure 5 shows activity coefficients determined by the vapor headspace technique for ethanol in the presence of varying amounts of water in both cyclohexanone and methylcyclohexanone. As shown in Table 4, the water-enhancement effect is only by factors of up to 1.5, for water:ethanol ratios up to 6.0 in two of the three cases. Also, there is no evidence that the activity coefficients level off at higher water contents. The degree of enhancement is much less than occurs for carboxylic acids and probably relates primarily to the activity coefficient of ethanol in water at high dilution being lower than those for ethanol in the cyclic ketones, i.e., a monotonic relationship. Diols. Figure 6 shows activity coefficients for 1,2propanediol (propylene glycol) in methylcyclohexanone with varying water contents, as measured by the vapor-
..
L
4'4
i
3.6 I ' 0
"
"
'
"
0.03 0.06 x
'.
I
"
"
0.09 0.12
"
"
0.15
[mol/mol]
Figure 6. Liquid-phase activity coefficients for 1,2-propanediol in methylcyclohexanone at 25 OC, as Influenced by the mole fraction of water in the liquid phase. 0.04 mol l,2-propanediol/mol methylcyclohexanone. Table 6. Solubilities of Phenols in Wet and Dry Solvents, 25 "C mole fraction solubility of solute solubility water (mole fraction) ratiob 0.04 0.82 0.44 0.42 0.51 2,3-xylenol MIBK 0.02 0.54 0.20 0.48 0.88 2,5-xylenol MIBK 0.06 0.51 0.20 0.47 0.91 2,6-xylenol MIBK 0.01 0.72 0.15 0.65 0.91 2,3-xylenol DIBK 0.0 0.49 0.04 0.48 0.98 2,5-xylenol DIBK 0.0 0.46 0.04 0.47 1.0 2-naphthol MIBK 0.01 0.19 0.17 0.16 0.9 2-naphthol CH 0.01 0.24 0.20 0.29 1.2 2-naphthol BA 0.01 0.28 0.18 0.20 0.7 solute phenol
solventa MIBK
a BA = n-butyl acetate, CH = cyclohexanone, DIBK = diisobutyl ketone, MIBK = methyl isobutyl ketone. Ratio of solute solubility at higher water content to solubility at lower water content.
headspace technique. The reduction in the activity coefficient from very low water content up to xw = 0.15 is by a factor of approximately 1.6. The activity coefficient may level off for n, > 0.10 (xw/xs> 2.51, suggesting that there may be specific ternary complexation or solvation. However, as for ethanol, the decrease in activity coefficient in the presence of water is probably simply a reflection of the activity coefficient of 1,2-propanediolin water at high dilution being lower than that for 1,2-propanediol in methylcyclohexanone. Vapor-headspace experiments were also carried out for 2,3-butanediol (mixed isomers) in methylcyclohexanone. There was substantial scatter in the data, with most of the computed activity coefficients ranging from 3.3 to 5.0 (Lee and King, 1993) and no single trend discernible in the data. Phenols. A number of different phenols were tested in methyl isobutyl ketone, cyclohexanone, and n-butyl acetate as solvents. Results are shown in Table 6. The solubilities of the phenols are all relatively high, as is also the case for the gallic acid results reported in Table 5. The high solubilities, even for 2-naphthol and gallic acid which have relatively high melting points, reflect the strong hydrogen bonding of the phenolic group with these electron-donor solvents. As a result of the high solubilities, the molar ratio of water to phenolic solute never exceeds approximately 1.0.
1378 Ind. Eng. Chem. Res., Vol. 33, No. 5, 1994
Also striking is the fact that solubilities are actually reduced by the presence of water except for 2-naphthol in cyclohexanoneand possibly also 2,5-xylenoland 2,3-xylenol in diisobutyl ketone. The lack of decrease in activity coefficient with added water probably primarily reflects the strong hydrogen bonding between the phenolic -OH groups and the electron-donor solvents, which results in very high equilibrium distribution coefficients of most phenols from water into ketones and esters (Earhart et al., 1977; Greminger et al., 1982); Le., the activity coefficients of phenols in ketones and esters are much lower than those in water. Therefore insertion of water in between the phenol and solvent molecules would not be favored energetically. The results shown for gallic acid in Table 5 therefore probably reflect enhanced solvation attributable to the carboxylic group, offset by reduction of the effect due to the strong, direct affinities of the phenolic -OH groups for the donor solvents. The enhancement of solubility observed for 2-naphthol in cyclohexanonemay reflect the greater steric accessibility of the carbonyl group in the cyclic ketone, as compared with the alkyl ketone. Conclusions Evaluation of data in the literature suggested that waterenhanced solvation could occur for carboxylic acids, alcohols, diols, and phenols as solutes in electron-donor solvents, such as ketones and esters. Two types of experiments were carried out with various of these solutes and solvents-measurements of the solubility of the solute in the solvent and vapor headspace analyses of the change in equilibrium partial pressure of the solute over the solvent, both as functions of the water content of the solvent. Results showed strong enhancement by water of the solvation of carboxylic acids in ketone and ester solvents. For cyclohexanone and methylcyclohexanone as solvents, activity coefficients decreased in the presence of water by factors of 2-3 for monocarboxylic acids (acetic and propionic), 6-8 for dicarboxylic acids (adipic, fumaric, and succinic), and 7-10 for a tricarboxylic acid (citric). These strong effects cannot be rationalized by the relative activity coefficients of carboxylic acids in otherwise pure water and otherwise pure solvent. They are probably attributable to strong ternary complexes or solvates, in which water takes a position between the carboxylic group and the electron-donor solvent. Similar measurements for ethanol and 1,Zpropanediol in cyclic ketone solvents showed decreases in activity coefficients of factors of 1.5-1.6,even in the presence of relatively large amounts of water. These effects are probably primarily associated with the fact that the activity coefficients in water are lower than those in ketones. Solubility measurements for phenol, several isomeric xylenols, and 2-naphthol in methyl isobutyl ketone and n-butyl acetate showed decreasesin solvation,i.e., increases in solute activity coefficients, due to water. This probably reflects the strong hydrogen bonding between phenols and these solvents, which results in activity coefficients of most phenols in ketones and esters being much lower than those in water. For gallic acid in cyclic ketones, addition of water to the solvent increased thesolute activity coefficient. The increase in solvation for gallic acid due to codissolved water is attributable to the carboxylic group, not the -OH groups. Water enhancement of solvation seems to be sufficient to be attractive as the basis for separation processes only in the case of carboxylic acid solutes, among the systems
studied. Examples of separations of this sort are precipitation of solute from a solvent caused by selective removal of codissolved water and enhancing leaching of carboxylic acids from sorbents by adding water to the solvent. Acknowledgment Financial support was provided by the Biochemical and Chemical Technology Research (BCTR) Program, Advanced Industrial Concepts Division, Office of Industrial Processes, Assistant Secretary for Energy Efficiency and Renewable Energy, U.S. Department of Energy, under Contract No. DE-AC03-76SF00098. The University of South Carolina provided sabbatical leave to one of the authors (Van Brunt) to enable him to carry out this work at the University of California, Berkeley and the Lawrence Berkeley Laboratory. Nomenclature E = solvation enhancement factor due to water [= -A(log Ye)/(XW/%) 1 W , = weight fraction of water in the liquid phase x, = mole fraction of the solute in liquid phase xw = mole fraction of water in the liquid phase ye = activity coefficient of solute in the liquid phase
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Received for review November 17, 1993 Accepted March 1, 1994. @
Abstractpublishedin Advance ACSAbstracts, April 1,1994.
