Langmuir 1990,6, 230-235
230
transition surface pressure a t low pH but has no effect at pH 1 5 . Within the condensed surface phase, there is little change in area, but there is some expansion in the Le state. The results are explained by penetration of the phosphotungstate anion between the polar head groups of the phospholipid, with the maximum effect corresponding with one anion to three phospholipid molecules. (5) UAc causes contraction of expanded phospholipid monolayers by lowering the transition surface pressure. The effect decreases with increasing pH, becoming insignificant a t pH 19. The effect is attributed to electrostatic attraction of adsorbed UOZ2+(or (U0,),(0H),")
for the negatively charged phosphate and the consequent drawing of the phospholipid head groups together. Clearly, the use of PTA r U A c as "stains" in preparing specimens containing phospholipids for examination in the electron microscope can lead to significant distortions if the pH is low.
Acknowledgment. We are indebted to Dr. M. L. Hair, Xerox Research Centre of Canada, Mississauga, for suggesting this problem, which was based on unpublished experimental work performed at the California Institute of Technology.
Solubilization of Mono- and Dichlorophenols by Hexadecylpyridinium Chloride Micelles. Effects of Substituent Groups Byung-Hwan Lee,t Sherril D. Christian,*pt Edwin E. Tucker,+and John F. Scamehord Institute for Applied Surfactant Research, T h e University of Oklahoma, Norman, Oklahoma 73019, Department of Chemistry and Biochemistry, T h e University of Oklahoma, Norman, Oklahoma 73019, and School of Chemical Engineering and Materials Science, T h e University of Oklahoma, Norman, Oklahoma 73019 Received April 4, 1989. I n Final Form: July 14, 1989 Semiequilibrium dialysis measurements have been made to determine solubilization equilibrium constants and activity coefficients of mono- and dichlorophenols in aqueous solutions of hexadecylpyridinium chloride (CPC) at 25 "C. Solubilization equilibrium constants (which relate the concentration of solubilizate in the micelle to the concentration of free organic in the "bulk solution") are well correlated by the expression K = Ko(l - BX)', where B is an empirical constant, KO is the limiting value of the solubilization constant, and X is the mole fraction of the phenol present in the surfactant micelles. Values of B are reasonably well linearly correlated with In KO,making it possible to predict entire solubilization isotherms from values of K obtained in the region of small X values. The effect of adding chloride groups is to increase the value of K at any value of X , although positional effects are also important. KOand B tend to be larger for compounds having chloro groups meta and para to the phenolic OH group.
Introduction Chlorinated aromatic compounds are common pollutants in ground water and aqueous industrial process streams.' Compounds such as mono- and dichlorophenols are soluble enough to be present in water a t concentrations of a t least 10-100 parts per million, so these substances can pose a very considerable threat to the environment and in particular to our water resources. Micellarenhanced ultrafiltration (MEUF) is a technique that can be used to remove a wide variety of organic pollutants, including chlorinated aromatic compounds, from In MEUF, surfactant is added to the aqueous stream, causing organic solutes to be solubilized within
' Department of Chemistry and Biochemistry.
* School of Chemical Engineering and Materials Science.
(1) Malaiyandi, M.; Wightman, R. H.; LaFerriere, C. Adu. Chem. Ser. 1987,214, 163. Matsuura, T.; Sourirajan, S.Adu. Chem. Ser. 1987, 214, 139. Robertson, J. H.; Cowen, W. F.; Longfield, J. Y. Chem. Eng. 1980, June 30, 102.
the micelles. The solution is then processed by ultrafiltration, using a membrane with pores small enough to reject the micelles and the solubilized organic; the permeate passing through the membrane consists of nearly pure water. Information about equilibrium solubilization of solutes in micelles can be applied directly to predict the results of MEUF separations of organic pollutants from There is little quantitative information about the solubilization of chlorinated phenols by aqueous surfactant micelles; hence, we have begun a series of studies of several classes of halogenated aromatic solutes. In the present investigation, the three monochlorophenols and all of the dichlorophenols except the 3,5 compound have been stud(2) Christian, S. D.; Scamehorn, J. F. In Surfactant-Based Separation Processes; Scamehorn, J . F., Harwell, J. H., Eds.; Marcel Dekker: New York, 1989; Chapter 1. (3) Dunn, R.0.; Scamehorn, J. F.; Christian, S. D. Sep. Sci. Technol. 1985, 20, 257; 1987, 22, 763.
0 1990 American Chemical Society
Substituent Group Effects on Solubilization ied as solubilized species in micelles of the cationic surfactant hexadecylpyridinium chloride (cetylpyridinium chloride or CPC). Of the literally hundreds of studies of the solubilization of organic compounds by surfactant micelles,4-14only a few have produced detailed and accurate solubilization isotherms, that is, measured solubilization results at varying solute activities and mole fractions of the solute in the micelles. Solubilization studies in this laboratory have provided accurate vapor pressure results for volatile hydrocarbons in anionic and cationic micelles.'"'' Numerous solubilization data have also been obtained for polar organic solutes by using the semiequilibrium dialysis (SED) meth~d'l-'~ and head-space chromatography.26 In SED experiments, ordinary equilibrium dialysis cells are used with membranes permeable to small molecules or ions (such as the organic solute and surfactant monomers) but impermeable to the surfactant micelles. The slow migration of surfactant through the membrane (over a period of 16-24 h) occurs simultaneously with the migration of the unsolubilized organic solute, which ordinarily diffuses rapidly enough to reach equilibrium with the solutions on both sides of the membrane. In calculating reliable values of solubilization equilibrium constants, one must allow for the presence of small concentrations of micellar surfactant in the permeate solution a t semiequilibrium; however, the corrections needed to account for this effect can be made with considerable accuracy, so the SED method can be used in solubilization studies with a wide range of organic solutes. SED is convenient to use in studying the solubilization of polar derivatives of aromatic hydrocarbons, because the concentrations of these substances can be determined, simultaneously with the surfactant, by UV spectral analysis. (4) Dougherty, S. J.; Berg, J. C. J. Colloid Interface Sci. 1974,48, 110. (5) Mukerjee, P. Solution Chemistry of Surfactants; Plenum Press: New York, 1979 Vol. 1, p 153. (6) Mukerjee, P.; Cardinal, J. R. J. Phys. Chem. 1978,82,1620. (7) Nagarajan, R.; Chaiko, M. A.; Ruckenstein, E. J. Phys. Chem. 1984,88,2916. (8) Bunton, C. A.; Sepulveda, L. J. Phys. Chem. 1979,83,680. (9) Hirose, C.; Sepulveda, L. J. Phys. Chem. 1981,85,3689. (10) Simon, S. A.; McDaniel, R. V.; McIntosh, T. J. J.Phys. Chem. 1982,86,1449. (11) Abuin, E. B.; Valenzuela, E.; Lissi, E. A. J. Colloid Interface Sci. 1984,101,401. (12) Goto, A,; Endo, F. J . Colloid Interface Sci. 1978,66, 26. (13) Moroi, Y.;Matuura, R. J . Colloid Interface Sci. 1988,125,463. (14) Valenzuela, E.; Abuin, E.; Lissi, E. A. J. Colloid Interface Sci. 1984,102,46. (15) Christian, S. D.; Tucker, E. E.; Lane, E. H. J. Colloid Interface Sci. 1981,84,423. (16) Tucker, E. E.; Christian, S. D. Faraday Symp. Chem. Soc. 1982. 17. 11. (17) 'Christian, S. D.; Smith, L. S.; Bushong, D. S.; Tucker, E. E. J . Colloid Interface Sci. 1982,89,514. (18) Tucker, E. E.; Christian, S. D. J. Colloid Interface Sci. 1985, 104,562. (19) Smith, G. A.; Christian, S. D.; Tucker, E. E.; Scamehorn, J. F. In Use of Ordered Media in Chemical Separations; Hinze, W. L., Armstrong, D. W., Eds.; ACS Symposium Series 342, American Chemical Society: Washington, DC, 1987;Chapter 10,p 184. (20) Smith, G.A.; Christian, S. D.; Tucker, E. E.; Scamehorn, J. F. J . Colloid Interface Sci. 1989,130,254. (21) Christian, S. D.; Smith, G. A.; Tucker, E. E.; Scamehorn, J. F. Langmuir 1985,1, 564. (22) Smith, G. A.; Christian, S. D.; Tucker, E. E.; Scamehorn, J. F. J. Solution Chem. 1986,15, 519. (23) Bhat, S.N.; Smith, G. A.; Tucker, E. E.; Christian, S. D.; Smith, W.; Scamehorn, J. F. Ind. Eng. Chem. Prod. Res. Deu. 1987,26, 1217. (24) Higazy, W.S.;Mahmoud, F. Z.; Taha, A. A.; Christian, S. D. J. Solution Chem. 1988,17,191. (25) Smith, G.A.; Christian, S.D.; Tucker, E. E.; Scamehorn, J. F. Langmuir 1987,3,598. Smith, G.A. Ph. D. Dissertation, The University of Oklahoma, Norman, 1986. (26) Nguyen, C.; Scamehorn, J. F.; Christian, S. D. Colloids Surf. 1988,30, 335.
Langmuir, Vol. 6, No. 1, 1990 231 In representing solubilization results, we adopt a viewpoint similar to that of others who have treated the micelle, containing both surfactant molecules and the solubilized organic substances, as a p~eudophase.'',~ We define the activity coefficient of an organic solute in the micelle (yA) as f A / ( f A o x ) , where f A is the fugacity of the solute, f A o is the fugacity of the solute in the standard state, and X is the mole fraction of the solute in the micelle. A pure component (preferably liquid) standard state is used, and X is defined on the basis of the number of moles of organic solute and number of moles of surfactant within the micelle, without considering the presence of water or electrolytes incorporated in or bound to the micelle. The solubilization equilibrium constant is a type of distribution coefficient, defined by K = X / C , , where cA is the concentration of monomeric or unsolubilized solute in the "bulk" aqueous solution. Because f A is proportional to cA in the very dilute aqueous solution of the organic solute investigated here, K varies inversely with Y~. Thus, polar compounds that have a strong tendency to solubilize within surfactant micelles are characterized by large values of K and (ordinarily) small values of yA. Although equilibrium solubilization data are intrinsically valuable, these results are also useful in indicating the locus of solubilization of molecules within surfactant micelles. Typically, polar organic solutes such as the aliphatic alcohols, carboxylic acids, phenols, and cresols have very small values of activity coefficients, and these values gradually increase as X increases. The solubilization results are consistent with other physical evidence indicating that these molecules have their head groups anchored in the polar/ionic outer region of typical ionic surfactants. The aliphatic or aromatic moieties of these polar solutes tend to solubilize a t least partly within the hydrocarbon core of the micelle, although steric and substituent group effects may play important roles in modifying the structure and thermodynamic properties of the "intramicellar solutions" of the organic solute in the micelles. Information about the activity coefficient of solubilized components and the dependence of yA on X may be useful in determining the environment of these compounds within the micelle. Values of yA near unity indicate that the environment of a solubilized species is energetically similar to that of the pure component (liquid or solid). Large values of yAaare characteristic of compounds that solubilize in environments less favorable than the pure solute state, and values of yA considerably less than unity indicate that the micelle strongly attracts the solute. Aliphatic hydrocarbons undoubtedly solubilize primarily within the hydrocarbon core region of micelles of the alkyl sulfates and CPC.15*20Solubilization isotherms for these very hydrophobic solutes typically exhibit activity coefficient vs X curves that decrease from relatively large values at infinite dilution to lower values as X increases toward unity. Although yA may be as large as 5 or 10 for alkanes in surfactant micelles a t small values of X , it should be emphasized that yA for these solutes in pure water is on the order of lo5 (the reciprocal of the mole fraction solubility of the hdyrocarbon in water). There has been considerable discussion in the literature about the location of aromatic solutes such as benzene and toluene in ionic surfactant micelles. Various types of physical evidence have been given to support claims that benzene solubilizes primarily in the surface (27) Shinoda, K.;Hutchinson, E. J . Phys. Chem. 1962,66,577. (28) McBain, M. E.; Hutchinson, E. Solubilization and Related Phenomena; Academic Press: New York, 1955;pp 138-143.
232 Langmuir, Vol. 6, No. 1, 1990
region of the m i ~ e l l eor, ~primarily ~ ~ ~ ~ within the micellar i n t e r i ~ r , ' ~ .or ' ~ in . ~ both ~ states.6 Extensive and precise solubilization results obtained by the vapor pressure method, for aromatic hydrocarbons in a variety of surfactants, do not indicate a strong preference for the solubilization of these compounds in either the micellar surface or hydrocarbon core r e g i ~ n . ~ The ~ , aromatic ~~ hydrocarbons are intermediate in behavior between highly polar solutes, which are clearly anchored in the micelle surface region, and aliphatic hydrocarbons, which preferentially solubilize in the hydrocarbon core region.l6Sz6 The activity coefficients of benzene and toluene are somewhat greater than unity in several anionic and cationic surfactants at small values of X,although these Values are considerably smaller than those for the aliphatic hydrocarbons. In surfactants like CPC and hexadecyltrimethylammonium chloride, the activity coefficients of benzene and toluene increase gradually from values near unity a t X = 0, reaching maxima near X = 0.5 or 0.6. A detailed thermodynamic analysis of vapor pressure results for the hydrocarbon/surfactant systems has been made in a number of cases, yielding virtually the only accurate information about free energies, entropies, and energies of transfer of solutes into micelles as functions of intramicellar c o m p o s i t i ~ n . ~ ~ , * ~ Several functional forms have been used to represent the dependence of the activity coefficient or the solubilization equilibrium constant on X.21-23For intramicellar solutions of polar organic solutes at X less than 0.2 or 0.3, it is often adequate to express K as a linear function of X,in the form K = Ko(l - b X ) . This leads to values of the solute activity coefficient that are proportional to 1/(1- b X ) , and a Gibbs-Duhem type integration may be used to infer values of the activity coefficient of the surfactant in the assumed micelle/solute pseudophase. However, we have found that if values of X as large as 0.3 to 0.6 are employed for typical polar solutes, a considerable upward curvature in plots of K vs X is observed, indicating that a quadratic or higher order term should be included in the function used to represent the dependence of K on X. The equation K = Ko(l - a X + b X 2 )reliably represents solubilization results for a wide variety of systems. Interestingly enough, a special form of this quadratic equation K = KO(1 - B X ) 2 (1) provides a n excellent fit of solubilization data for phenol and for all of the mono- and dichlorophenols in aqueous CPC solutions, with no significant loss of precision in fitting results, as compared to the general quadratic expression, K = Ko(l - a X + b X 2 ) . We believe that eq 1, which readily yields expressions for the activity coefficients of both the organic solute and the surfactant within the micelle, will be convenient for fitting a variety of solubilization results, throughout wide ranges of intramicellar compositions. (29) Eriksson, J. C.; Gillberg, G. Acta Chem. Scand. 1966,20,2019. (30) Grieser, F.; Drummond, C. J . J . Phys. Chem. 1988, 92, 5580. (31) Rehfeld, S. J . J . Phys. Chem. 1971, 75, 3905; 1970, 74, 117. (32) Christian, S. D.; Tucker, E. E. Am. Lab. (Fairfield, Conn.) 1982. 24(8). 36. (33) Christian, S. D.; Tucker, E. E. Am. Lab. (Fairfield, Conn.) 1982, 14(9),31. (34) Tanford, C . J. Phys. Chem. 1972, 76, 3020. (35) Tartar, H. V. J . Colloid Sci. 1959, 14, 115. (36) Nakagawa, T.; Kuriyama, K.; Inoue, H. J. Colloid Sci. 1960, 15., 268. ~ . . ~
(37) Reiss-Husson, F.; Luzzati, V. J. Phys. Chem. 1964, 68, 3504. (38) Ekwall, P.; Mandell, L.; Solyom, P. J . Colloid Interface Sci. 1971, 35, 519. (39) Hansen, R. S.; Miller, F.A. J . Phys. Chem. 1954,58, 193. (40) Nagakura, S. J . Am. Chem. SOC.1954, 76, 3070.
Lee et al.
Experimental Section The semiequilibrium (SED)method has been extensivelyused in our laboratory to study the solubilization of organic and ionic species by surfactant micelles in aqueous solutions. The method utilizes ordinary equilibrium dialysis cells (Fisher Scientific Co.), in which a retentate solution, containing a concentrated surfactant solution and a dissolved organic or ionic solute, is placed on one side of the membrane and a permeate solution, consisting of pure water or a chosen electrolyte solution, is placed on the other. Ordinarily, 6000 dalton molecular weight cutoff membranes (regenerated cellulose),carefully washed, are used to separate the solutions and to prevent transfer of micelles from the retentate to the permeate. At semiequilibrium, usually attained within 16-20 h, the solute species has attained equilibrium with the surfactant solutions on both sides of the membrane; that is, the concentration of organic solute in the permeate solution will be approximately equal to cA, neglecting the amount of the solute in the micelle in the permeate compartment. Although equilibrium is readily reached by the organic solute, the surfactant continues to transfer slowly from the retentate to the permeate solution. Typically, the concentration of surfactant in the permeate at semiequilibrium is 10-30% greater than the critical micelle concentration. In the studies described here, the concentration of hexadecylpyridinium chloride (CPC) is varied from 0.025 to 0.100 M, and the chlorinated phenols are introduced at mole ratios (solute to surfactant) varying from about 0.1 to 1.0. The surfactant concentration in the retentate remains large compared to M). In contrast the critical micelle concentration (8.8 X to some previous studies using the SED method, analyses were made of both the surfactant and the organic solute in the two compartments at semiequilibrium. Osmotic pressure effects are large enough to cause a considerable volume of water to transfer from the permeate to the retentate solution. In the present experiments, we observed 10-40% increases in the retentate volumes (and correspondingdecreases in permeate volumes) after about 20 h. If one were to estimate the concentrations of surfactant and organic solute in the retentate at equilibrium, by difference from analyses of the permeate solution only, erroneous values would be inferred if corrections were not introduced to account for the transfer of water. Interestingly, values of solubilization equilibrium constants inferred by analyzing only the permeate solution (and neglecting the volume change) are nearly correct, because compensating errors are made in the calculated concentrations of both the organic and the surfactant. However in the most careful work it is necessary to determine the concentrations of the organic solute and the surfactant in both compartments. Aldrich Gold-label (99%) phenol and the monochlorophenols, dichlorophenols (2,3, 2,4, 2,5, 2,6 and 3,4), and hexadecylpyridinium chloride (CPC) from Hexcel were used as received. The CPC does not exhibit a minimum in a plot of surface tension vs concentration. Double-distilled,deionized water was used in all experiments. The permeate and retentate solutions were analyzed for CPC and the aromatic solute by multiwavelength UV spectroscopy. Absorbances were measured at six wavelengths chosen near the absorption maxima for the respective solutes, and concentrations were inferred by linear least-squares analysis, using known absorptivities of the solutes. Data Analysis Several models have been developed to relate solubilization constants for polar solutes to the mole fraction (X)of the solute in surfactant micelles (vide supra). The simple equation K = Ko(l - BX)' provides an excellent fit of data for the systems studied here and for that matter for almost all of the solubilization data for CPC obtained previously with the SED method. Some numerical analysis of the data is required to infer values of K from the measured concentrations of surfactant and solute in the two compartments of the SED cell. In order to relate the total concentration of organic solute (A) and CPC in either compartment to X and Val-
Langmuir, Vol. 6, No. 1, 1990 233
Substituent Group Effects on Solubilization ues of the activity coefficients, one can use the equations
[A],, = YAXCA"+ (X/(1- X))[CPClmic
(2)
(3) [CPC],, = ~ c p c ( 1 - X ) c c p c " + [CPClmic where the subscripts tot and mic refer to total and intramicellar concentrations, respectively. cA0 is a limiting concentration of A consistent with the pure-component standard state (see next paragraph), cFpc0is the concentration of monomeric CPC in water in the absence of added organic (equal to the critical micelle concentration), and y A and ycpcare activity coefficients of the organic solute and the surfactant, based on the pure component and pure micelle standard states, re~pectively.''~'~ If the activity coefficients of both components in the micelle were known explicitly as functions of X, it would be possible to solve eq 2 and 3 simultaneously to calculate X and [CPCImicfor both the permeate and retentate solutions. Previously, we expressed the activity coefficients in forms suggested by Hansen and Miller3' and also in forms consistent with an assumed linear dependence of K on X. It is also possible to derive explicit relations for the activity coefficients consistent with eq 1. For a sparingly soluble organic solute, the activity of the solute in its saturated aqueous solution will be the same as that of the pure organic compound, provided the pure compound does not dissolve an appreciable concentration of water. Therefore, the activity of the organic solute in the micelle, on the pure-component standard state basis, is given by aA = c A / c A " , where cA is the concentration of monomer in the aqueous solution and cA0 is the saturation concentration of the organic in water (in the absence of surfactant). (If the solute is not sparingly soluble in water, or if the organic compound dissolves appreciable amounts of water, the activity is still expressed as cA/cAo, but cA0 is now interpreted as the hypothetical concentration of monomeric organic solute (in aqueous solution) at which the solute activity will equal that of the pure organic component, for an assumed Henry's law behavior at concentrations up to that concentration.) In previous work, we have used pure component liquid or super-cooled liquid standard states, even for s o l u t e s t h a t a r e crystalline solids a t room temperat~re.'~-'~ However, in the present research, most of the organic compounds are solids for which the relative activities of solid and liquid a t 25 "C are not available in the literature. Therefore, we have used pure component (solid) standard states for the crystalline chlorinated phenols. Given that y A = a A / X and K = x / c A , and assuming the validity of eq 1, we can express the activity coefficient of the organic solute in the micelle as = l/(KCAo)= a / ( l - B X ) 2
(4) where a = l/KocAo. By utilizing eq 4 and the GibbsDuhem equation, we can obtain the equation YA
In ycpc= (2/(1- B ) ) [ BIn (1- X) - In (1- BX)] (5) to represent the dependence of the activity coefficient of CPC in the micelle on X . In the limit as X 0, ycpc1
-
1.
A nonlinear least-squares method described p r e v i ~ u s l y ~is' * utilized ~~ to analyze SED data for each system, consistent with the mathematical model described by eq 1-5. First, eq 2 and 3 are solved simultaneously, using analytical concentration data for each of the retentate solutions, to obtain approximate values of X and
Table I. Least-Squares Parameters for Mono- and Dichlorophenols in CPC at 25 "C compd phenol o-chloro m-chloro p-chloro 2,3-dichloro 2,4-dichloro 2,5-dichloro 2,6-dichloro 3,4-dichloro
KO,M-' 81 f 469 f 537 f 786 f 2243 f 2986 f 1908 f 727 f 3275 f
a
Bb
cAD,M'
1 21 25 49 96 128 47 20 165
1.09 f 0.01 1.28 f 0.02 1.25 f 0.02 1.25 f 0.03 1.26 f 0.01 1.38 f 0.01 1.34 f 0.01 1.21 f 0.01 1.39 f 0.02
0.8 0.1731 0.1408 0.2077 0.0252 0.0338 0.0228 0.0125 0.0530
lo'
RMSD,Md 2.108 1.547 1.529 1.507 1.459 0.524 0.463 1.031 0.490
"Intercept of a plot of the solubilization constant K vs the mole fraction of organic solute in the micelle, X. Parameter in eq 1. Constant in eq 4, defining activity coefficient of organic solute. dRoot mean square deviation in organic solute concentration in the permeate solution, fitted with model described.
[CPCImic. This is accomplished by using trial values of the variable parameters ( B and a), known values of the standard concentrations (cA0 and ccpco),and eq 4 and 5 to relate the activity coefficients to the trial parameter values and X. Next, eq 3 is employed to calculate [CPCIdc for each of the permeate solutions, assuming that X in the permeate solution in each experiment is the same as X in the retentate solution and using eq 4 and 5 with the same values of B and a to predict values of the activity coefficients. (The assumption that X is the same in the permeate compartment as in the retentate has been justified previously for several types of solutes in SED experiments.) Finally, the predicted concentration of organic solute in the permeate, [A]totcdcd,is inferred from eq 2 for each of the experiments. $he sum of squares of deviations between calculated and experimental values of [A],,, is in this way computed for all of the data sets for a given system. The nonlinear least-squares program varies B and a in an unrestricted way to obtain the minimum sum of squares of deviations and the optimum Values of the parameters. Table I contains the least-squares values of KOand B obtained by fitting data for phenol and the chlorinated phenols to the model described above. (Note that eq 4 implies that KO= l/acAo.) Also included in Table I are the values of cA0,approximately equal to the solubility of each compound, used in the model. The final column in the table lists values of the root mean square deviation (RMSD) in concentration of the organic solute obtained from the analysis. When the SED results for each system are examined, it is convenient to infer individual values of K, corresponding to measurements of the concentrations of the organic solute and the surfactant in both compartments of the equilibrium dialysis cell. Previously, we showed that K can be calculated from the equation
K = (1- X)([Alret,tot - [Al,r,,t)/([Al,r,,t[CPClr,t,mic
-
[Alr,t,,t[CPCl,r,mic) ( 6 ) where the subscripts per and ret refer to the permeate and retentate compartments and the subscript tot refers to total concentration.'l In eq 6, both [A],,, and [A],,, are directly measurable, and [CPC]ret,micand X can be inferred (for the retentate solution) by solving eq 2 and 3 simultaneously, with eq 5 (with the least-squares value of B ) being used to relate y A and ycpcto X . (Recall that X is assumed to be the same in the permeate solution as in the retentate.) Finally, [CPC]per,mic can be calculated by using eq 3, with the measured value of [CPC] er,tot and the value of ycpccalculated from X and the Peast-
234
Langmuir, Vol. 6, No. I , 1990 30
Lee et al.
I 0.10-
,n
: 20
0 08
i:
4
>
0060.04
10
-
0.020
0
01
02
03
04
05
06
.
I
01
.
,
02
.
I
03
.
,
04
.
,
05
. 06
X
X
solute mole fraction in micelle Figure 1. Dependence of solubilization equilibrium constants for monochlorophenolsand phenol in hexadecylpyridinium chloride micelles on intramicellar mole fraction of solute. 50
.
I
I
I
solute mole fraction in micelle Figure 3. Activity coefficients of monochlorophenols and phenol in hexadecylpyridinium chloride micelles as a function of intramicellar mole fraction of solute. surface sites. It is possible to rearrange eq 7 in the form which may be compared with the Langmuir equation written as
0 2
0 3
05
0 4
0 6
0 7
X
solute mole fraction in micelle Figure 2. Dependence of solubilization equilibrium constants for dichlorophenols in hexadecylpyridinium chloride micelles on intramicellar mole fraction of solute. squares value of B (using eq 5 ) . Thus, individual K values can be calculated from eq 6 from the results of each dialysis experiment, once the parameter B in eq 5 has been inferred from the least-squares analysis of all the data sets for a given system.
Results and Discussion The model described above implies that a plot of the square root of K against X for each system will be linear, with a slope equal to -B. Figures 1 and 2 show results plotted in this form; the K values have been calculated for the individual experiments by using the method described in the preceding paragraph, and the lines are drawn to correspond to the least-squares values of B and KOlisted in Table I. Excellent conformity to the model is observed for all of the data for phenol and the chlorinated phenols. In the case of the dichlorinated phenols, values of X plotted in Figure 2 (and fitted with the model) have been limited to 0.2 or greater. Relative errors in the chemical analyses become significant for the SED experiments a t X < 0.2, because only very small concentrations of the strongly solubilized dichlorophenols are able to diffuse into the permeate solution. Although K depends quadratically on 1 - B X , the dependence of K on X a t small values of X is nearly linear. Thus, expanding eq 1 and ignoring the term in X 2 , we can write K = K,(1 - 2 A X ) , which is consistent with Langmuir's adsorption e q ~ a t i o nin~ the ' ~ ~form
If the adsorption of the polar solutes occurs with the phenolic OH group bound in the head-group region of the micelle, an increase in the amount of adsorption of the phenols as cA increases may be considered to "use up"
where 8 is the fraction of the surface "sites" that are occupied by bound phenol molecules and a is the concentration of the phenol in the bulk aqueous solution a t which B = 0.5. Thus, 2BX is equal to 0, which indicates that the ratio of the number of surfactant molecules to the number of phenol molecules per site is equal to 2B. For all of the chlorinated phenols studied here, 2B is in the range 2-3, and B increases gradually with the strength of the interaction between the phenols and the CPC micelles. One interpretation of this observation is that each phenol molecule may interact strongly with two or three CPC molecules, so these molecules can no longer act as a primary solubilization site for additional organic solute molecules. It is not unreasonable to expect that the packing of the pyridinium groups in the micellar surface region would allow as many as three such groups to interact strongly with the OH group of a single phenol molecule. By using the extended-chain model for an assumed spherical micelle, one can estimate that the average area of each pyridinium head group in CPC is approximately 60 A2.34*35Thus, the area per phenol molecule corresponding to a "site" is approximately 120-180 A2 (2-3 times the average area per head group). Although it might be interesting to speculate about the nature of sites occupied by alcohols, phenols, and related polar solutes bound to ionic surfactant micelles, the interpretation probably should not be carried too far, because the Langmuir adsorption model only applies a t relatively small values of X . Eventually, the binding of a sufficient number of polar molecules in the surface region may be expected to drastically alter the size and shape of ionic micelle^.^^-^* Figures 3 and 4 show the activity coefficients of phenol and the chlorinated phenols plotted against X,with the pure solid standard states being used for all of the compounds except phenol, o-chlorophenol, and p chlorophenol. All of these activity coefficient plots show the expected behavior for polar solutes interacting strongly with the polar ionic head groups of CPC. At low concentrations of the solute, activity coefficients on the order of 0.1 or less are observed; as X increases, values of yA increase gradually toward unity, although the curves must terminate as the limiting solubility of the organic solute in t h e micelle is reached. Activity coefficients of the sur-
235
Substituent Group Effects on Solubilization 1.0
9-
/ I
I
I
3,4-dichloro 8-
0.6
2.3-dichloro
d
-
4
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0.2
0.3
0.4
0.6
0.6
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.
,
.
,
.
0.7
X
solute mole fraction in micelle Figure 4. Activity coefficients of dichlorophenols in hexadecylpyridinium chloride micelles as a function of intramicellar mole fraction of solute.
factant, on the basis of the pure micelle standard state, are not included in the plots, but these may be readily calculated for each system by using eq 5; only the value of the parameter B is required to calculate ycpcas a function of X. The activity coefficient vs X curve for 2,6-dichloropheno1 is displaced toward considerably larger values of the activity coefficient than that for the other dichlorophenols, consistent with the observation (Table I) that the value of KO for 2,6-dichlorophenol is much smaller than that for the other dichloro compounds. It is probable that the presence of two chloro groups ortho to the phenolic OH group greatly modifies the interaction of this group with the pyridinium moieties of the CPC molecules. Intramolecular hydrogen bonding is thought to occur between the phenolic OH group and ortho chloro groups.40 It should also be noted that 2-chlorophenol has a smaller KOvalue than 3- or 4-chlorophenol. In general, the results in Table I indicate that the further away from the OH group and C1 is located the greater will be the enhancement in KOcaused by adding the group. Chloro groups act similarly to the hydrophobic methyl groups in enhancing the solubilization of aromatic compounds, and in fact a chlorosubstituent is even more effective than the methyl group in increasing the value of KW In the case of methyl groups, positional effects seem to be less important; KOvalues for o-, m-, and p-cresol have been reported to be 187, 190, and 195 M-l, respect i ~ e l y . ' ~In the case of C1, electronic effects as well as hydrophobic and steric effects are involved; thus, the acid strengths of phenols are considerably increased by substituent C1 groups, particularly at the ortho position. The inductive effect of C1-which increases the acidity of the OH proton and decreases the basicity of the oxygenprobably contributes to a decrease in the strength of the interaction between the polar OH group and the positive pyridinium group. However, the interaction of the OH group of phenol with the counterion (C1-) should be strengthened by the presence of C1 on the aromatic ring near the OH group. We have assumed that the phenols are solubilized primarily in the neutral form because of the solution pH (ca. 4.2-5.2) and the fact that titration of a 3,4-dichlorophenol (0.01 M)-CPC (0.05 M) mixture with NaOH leads to precipitate formation whereas titration of CPC alone does not. Additional studies with other electron-withdrawing and electron-donating groups and with other surfactants should help determine the relative importance of inductive, steric, and hydrophobic effects in stabilizing solubilized polar molecules. An examination of the KOand B values in Table I shows that these values are reasonably well correlated. Thus, B is smallest for phenol, which has the smallest value of
Figure 5. Dependence of limiting solubilization equilibrium constant (KO)on values of the parameter B in eq 1.
KO,and B tends to increase as KOincreases. A plot of In KO vs B (Figure 5 ) shows the correlation, which is reasonably good, although the points for three of the systems (which contain C1 at the ortho position and in which OH4!1 hydrogen bonding probably occurs) deviate appreciably from the fitted curve. The equation of the leastsquares straight line (obtained by fitting data for all of the systems except o-chlorophenol and 2,3- and 2,6dichlorophenol) is In KO = -9.3 + 12.6B. Combining this relation with eq 1, one can write
K =.(l- B X ) 2 exp(12.6B - 9.3) (10) to correlate most of the results obtained in the present study. Thus, to the extent that eq 10 can be trusted, it permits fitting results for individual solutes solubilized by CPC, throughout wide ranges of X,with a single value of the parameter B. In attempting to obtain data for the solubilization of chlorinated phenols, an important class of industrial and environmental pollutants,' the convenient form of eq 1 and 10 should simplify the task of obtaining the necessary experimental results to plan colloidenhanced separation processes. Finally, it should be mentioned that the values of K for the chlorinated phenols are large enough so that reasonable concentrations of surfactant can be used to solubilize most of these compounds in aqueous streams. If K is on the order of 2000 M-l (as it is for most of the dichlorophenols) and the surfactant concentration is 0.05 M, approximately 99% of the organic solute resides within the micelles. Thus, in MEUF separations2v3it should be feasible to achieve 2 order of magnitude reductions in concentrations of the dichlorophenols in a single ultrafiltration stage. In the case of the monochlorophenols, a 96-97% reduction in concentration can be achieved under analogous conditions. Acknowledgment. We appreciate the financial support of the Office of Basic Energy Sciences, Department of Energy, Contract DE-FG05-87ER13678, Department of Energy Grant No. DE-FG01-87FE61146,and National Science Foundation grant CHE 8701887. In addition, we gratefully acknowledge the assistance of industrial sponsors of the Institute for Applied Surfactant Research, including the Aqualon Co., Kerr-McGee Corp., Sandoz Chemicals Corp., E. I. du Pont de Nemours & Co., Unilever, Inc., and Shell Development Co. Registry No. CPC, 123-03-5;phenol, 108-95-2;o-chlorophenol, 95-57-8; m-chlorophenol, 108-43-0;p-chlorophenol, 106-489; 2,3-dichlorophenol, 576-24-9;2,4-dichlorophenol,120-83-2; 2,5dichlorophenol, 583-78-8; 2,6-dichlorophenol, 87-65-0; 3,4dichlorophenol, 95-77-2.
