Langmuir 1991, 7, 95-100
95
Solubilization of 2-Phenylethanol by Dodecyldimethylamine Oxide in Aqueous Solution Hirotaka Uchiyama,tJ Sherril D. Christian,*ptJ John F. Scamehorn,tJ Masahiko Abe,*.ll and Keizo Ogino*J Institute for Applied Surfactant Research, School of Chemical Engineering and Materials Science, and Department of Chemistry, The University of Oklahoma, Norman, Oklahoma 73019, Faculty of Science and Technology, Science University of Tokyo, Yamazaki, Noda, Chiba 270-01, Japan, and Institute of Colloid and Interface Science, Science University of Tokyo, 1-3, Kagurazaka, Shinjuku-ku, Tokyo 162, Japan Received April 5, 1990 The solubilization of 2-phenylethanol (PEA) was measured over a wide range of solute activities and pH values by using dodecyldimethylamineoxide (DDAO) as the surfactant. The DDAO in micellar form is all cationic (protonated) at low pH and all nonionic at high pH. At intermediate pH levels, the DDAO forms mixed micelles containing both the cationic and nonionic forms of surfactant. Thus, measurement of the solubilization of PEA as a function of pH produces solubilization data as a function of mixed micelle composition for this amphoteric surfactant. The solubilizationequilibrium constant was found to decrease with increasing mole fraction of PEA in the micelle for all pH values and to be less in the mixed micelles than in either pure cationic or pure nonionic micelles. This latter effect could be due to the hydrophilic region of the mixed micelle being more compact than that of the single-componentmicelles. Introduction Surfactants which are used in practical applications are almost always mixtures of surface-active compounds. Understanding how surfactants interact in mixed micelles is thus essential for the many industrial applications of surfactants. A number of solution properties have been measured for various mixed surfactant systems, including critical micelle concentration,'Y2 sol~bilization,~+j adsorpt i ~ n , precipitation!JO ~*~ and cloud points.11J2 In these papers, it has been reported that the properties of a mixed surfactant system are often superior to that of a singlesurfactant system. Dodecyldimethylamine oxide (DDAO) can be modeled thermodynamically as a binary mixture of cationic and nonionic surfactants, because the protonated and neutral ~~
* T o whom all correspondence should be addressed. t Institute for Applied Surfactant Research, The University of Oklahoma. J School of Chemical Engineering and Materials Science, The University of Oklahoma. f Department of Chemistry, T h e University of Oklahoma. I Faculty of Science and Technology, Science University of Tokyo. 11 Institute of Colloid and Interface Science, Science University of Tokyo. (1)Rosen, M. J.; Hua, X. Y. J. Am. Oil Chem. SOC.1982,59, 582. (2) Holland, P. M. Adu. Colloid Interface Sci. 1986, 26, 111. (3) Smith, G. A.; Christian, S. D.; Tucker, E. E.; Scamehorn, J. F. J. Colloid Interface Sci. 1989, 130, 254. (4) Nguyen, C. M.; Scamehorn, J. F.; Christian, S. D. Colloids Surf.
1988, 30, 335.
(5)Abe, M.; Kubota, T.; Uchiyama, H.; Ogino, K. Colloid Polym. Sci.
1989,267, 365. (6) Uchiyama, H.; Tukuoka, Y.; Abe, M.; Ogino, K. J.ColloidInterface Sci. 1989, 132, 88. (7) Scamehorn, J. F.; Schechter,R. S.; Wade, W. H.J. ColloidInterface Sci. 1982, 85, 494.
(8)Lopata, J. J.; Harwell, J. H.; Scamehorn, J. F. In Surfactant-Based Mobility Control;Smith, D. H., Ed.; ACS Symp. Ser.; American Chemical Society: Washington, D.C., 1988; Vol. 373, p 205. (9) Stellner, K. L.; Scamehorn, J. F. Langmurr 1989,5, 77. (10) Stellner, K. L.; Amante, J. C.; Scamehorn, J. F.; Harwell, J. H. J. Colloid Interface Sci. 1988, 123, 186. (11) Yoesting, 0. E.; Scamehorn, J. F. Colloid Polym. Sci. 1986,264, 148. (12) Maclay, W. N. J. Colloid Sci. 1956, 11, 272.
0743-7463/91/2407-0095$02.50/0
species of the DDAO molecule can be treated as separate surfactants, the composition of which is varied by adjusting the solution pH. The effect of pH and added electrolyte on the micelle formation properties of amine oxides13-18 and mixtures of amine oxide with other has been studied. One of the most important properties of micellar systems is their ability to solubilize organic species. Therefore, solubilization by surfactants is important in many fields such as enhanced oil recover^,^^'^^ cosmetics,26 micellar catalysis,27and surfactant-basedseparations.2g30 The solubilization equilibrium constants of organic solutes in surfactant solutions have been measured by various methods; (13) Hermann, K. W. J. Phys. Chem. 1962,66, 295. (14) Tokiwa, F.; Ohki, K. J. Phys. Chem. 1966, 70, 3437. (15) Ikeda, S.;Tsunoda, M.; Maeda, H. J. Colloid Interface Sci. 1978, 67, 336. (16) Ikeda, S.; Tsunoda, M.; Maeda, H. J. Colloid Interface Sci. 1979, 70, 448. (17) Imae, T.; Ikeda, S. Colloid Polym. Sci. 1985, 263, 756. (18)Chang, D. L.; Rosano, H. L.; Woodward, A. E. Langmuir 1985,1, 669. (19) Imae, T.; Araki, H.; Ikeda, S. Colloids Surf. 1986, 17, 207. (20) Imae, T.; Araki, H.; Ikeda, S. Colloids Surf. 1986,17, 221. (21) Kolp, D. G.; Laughlin, R. G ; Krause, F. P.; Zimmerer, R. E. J. Phys. Chem. 1963,67, 51. (22) Rosen, M. J.; Friedman, D.; Gross, M. J.Phys. Chem. 1964,68, 3219. (23) Chang, D. L.; Rosano, H. L. In StructurelPerformane Relation-
ships in Surfactants; Rosen, M. J., Ed.; ACS Symp. Ser.; American Chemical Society: Washington, D.C., 1984; Vol. 253, p 129. (24) Bourrel, M.; Schechter, R. S. Microemulsions and Related System-Formulation, Solvency, and Physical Properties; Marcel Dekker: New York, 1988; Chapter 7. (25) Miller, C. A.; Qutubuddin, S. In Interfacial Phenomena in Nonaqueous Media; Eicke, H . E.; Parfitt, G. D., Eds.; Marcel Dekker: New York, 1986; Chapter 4. (26) Lin, T. J. In Surfactants in Cosmetics;Rieger, M. M.; Ed.; Marcel Dekker: New York, 1985; Chapter 2. (27) Fendler, J. H.; Fendler, E. J. Catalysis in Micellar and Macromolecular Systems; Academic Press: New York, 1975; Chapter 4. (28) Scamehorn, J. F.; Harwell, J. H. Surfactant-based Separation Processes; Marcel Dekker: New York, 1989. (29) Dunn, R. 0.;Scamehorn, J. F.; Christian, S.D. Sep. Sei. Technol. 1985, 20, 257. (30) Dunn, R. 0.; Scamehorn, J. F.; Christian, S. D. Sep. Sci. Technol. 1987, 22, 763.
0 1991 American Chemical Society
Uehiyama et al,
96 Langmuir, Vol. 7, No. 1, 1991
see ref 31 for a review of experimental methods of measuring solubilization. In this study, the solubilization equilibrium constant has been determined by the semiequilibrium dialysis (SED) method, which has been extensively used in our laborat0ry.3~-3~ Perfume compounds, which are usually oil-soluble materials, are often used in practical applications as solutes solubilized by surfactants. However, there have been only a few papers on the solubilization of perfume compounds by various Therefore, we have investigated the solubilization of 2-phenylethanol, a commonly used compound in perfumes, by DDAO in aqueous solution. Experimental Section Materials, The dodecyldimethylamine oxide (DDAO) was obtained from Fluka Chem. Co. as a 30% aqueous solution. The water was removed by freeze-drying,and DDAO was recrystallized several times from acetone. This hygroscopic solid was dried and stored in vacuo. This purity was ascertained by surface tension measurements. 2-Phenylethanol(Gold label;PEA)from Aldrich was used as received. The phenolphthalein and &cyclodextrin were obtained from Aldrich,the sodium chloridefrom Mallinckrodt, the sodium carbonate from Fisher, and the hydrochloricacid from Curtin Matheson;all of these compounds were used as received. The water was distilled twice and deionized. Measurements. The semiequilibrium dialysis (SED) method3sM was used to measure solubilizationhere. This method utilizes ordinary equilibrium dialysis cells (FisherScientificCo.) with regenerated cellulose acetate membranes having a 6000dalton molecular weight cutoff. The dialysis membranes were washed carefully in distilled water for approximately 10h before use and placed between the compartments to separate the solutions to prevent transfer of micelles from one side to the other. In all experiments, 0.1 M NaCl was used as a swamping electrolyte to minimize the change in the ionic strength of the solution with varying pH. A 0.1 M NaCl solution containing known concentrationsof the DDAO and PEA was placed on one side of the membrane (retentate), and the other side (permeate) contained 0.1 M NaCl solution. The cells were thermostated at 25 "C for 24 h. The concentrationof DDAO in the permeate was determined by employing a visible spectral displacement method,a in which phenolphthalein is displaced from the 0-cyclodextrin cavity by the hydrocarbon moiety of the surfactant. The 1:l complex of phenolphthalein with 6-cyclodextrinhas no absorbance in the region of the 550-nm band of the basic form of phenolphthalein at a pH of approximately 10.5. Therefore, the concentration of the surfactant was determined directly from the absorbance at this wavelength. The concentration of 2-phenylethanol in the permeate was determined by UV spectroscopy (Varian Cary Model 118). The retentate PEA concentrationwas obtained by a material balance which accounts for the solutes transferred into the permeate compartment. The pH was monitored by using a Fisher Model 620 pH meter. After an initial calibration of the pH electrode, the amine oxide (31) Nguyen, C. M.; Scamehorn, J. F.; Christian, S.D. Tenside Surfactant Deterg. 1988, 25, 328. (32) Christian, S. D.; Smith, G. A.; Tucker, E. E.; Scamehorn, J. F. Langmuir 1985, 1,564. (33) Smith, G. A.; Christian, S.D.; Tucker, E. E.; Scamehorn, J. F. J. Solution Chem. 1986, 15, 519. (34,) Higazy, W. S.;Mahmoud, F. Z.; Taha, A. A.; Christian, S.D. J. Solutron Chem. 1988, 17, 191. (35) Bhat, S.N.; Smith, G. A,; Tucker, E. E.; Christian, S.D.; Scamehorn, J. F., Smith, W. Ind. Eng. Chem. Res. 1987,26, 1217. (36) Lee,B. H.; Christian,S. D.;Tucker, E. E.; Scamehorn, J. F.Langmuir 1990,6, 230. (37) Akahoshi, R.; Horike, S.;Noda, S. Nippon Kagaku Kaishi 1984, 1974. (38) Akahoshi, R.; Horike, S.;Noda, S.Nihon Kagaku Kaishi 1985, 943. (39) Saeaki, K. J.; Christian, S. D.; Tucker, E. E. J . Colloid Interface Sci. 1990, 134, 412.
solutions were titrated with HCl to approximately pH 3.5. The electrode was then recalibrated, using standards at lower pH. The titration was then continued down to a pH of 2.0.
The hydrodynamic size of the DDAO micelle was determined
a Submicron particle analyzer (Malvern Instrument, U.K., Model 4700). The optical source on the light-scatteringapparatus was an argon ion laser operating at 488 nm with an output power of 5 W maximum (Coherent Co., Model Innova 90). In order to get the mutual diffusion coefficient, the measurement of the timedependent correlation function of the scattered intensity at a scattering angle 90" was performed to determine the change of the micellar size with pH. The data analysis was performed by
the combined use of the cumulant method40 and the model free algorithm.40 The hydrodynamic radius was calculated by the Einstein-Stokes equation with the mutual diffusion coefficient. Before measurement, the aqueous solution of DDAO was passed 3 times through the membrane filter with a 0.1-pm pore size (Cellulosenitrate type of Toyo Roshi Co. Ltd. Tokyo) for optical purification.
Data Analysis The solubilization equilibrium constant of an organic solute in an aqueous micellar solution ( K ) is defined as the ratio of the mole fraction of organic solute in the micellar phase to the molar concentration of the unsolubilized monomeric organic solute in the aqueous phase. The mole fraction of organic solute in the micellar phase is defined as the molar concentration of the solute in the micellar phase divided by the total molar concentration of surfactant and organic solute in the micellar phase. The analysis that follows was developed for aqueous systems containing only one surfactant solute and a solubilized organic component. The approach is based on the pseudo phase equilibrium model, which treats the surfactant and dissolved organic solute in the micelle as being analogous to a separate condensed phase, existing in thermodynamic equilibrium with monomers of both the organic solute and the surfactant in the "bulk" aqueous phase. In the present study, the micelle actually contains three components: the surfactant in two distinct forms (neutral DDAO and the cationicDDAOH+) and the organic solute. However, at any given pH value, the mole ratio of DDAO to DDAOH+ does not change. Therefore, it seems reasonable to use the same phase equilibrium model a t each pH, again treating the "intramicellar solution" as a pseudo-two-component phase containing the surfactant and the organic solute. With this assumption, all of the equations developed p r e v i o u ~ l ycan ~ ~ be - ~ applied directly to infer activity coefficients and compositions of the intramicellar solutions in the permeate and retentate compartments of the SED cell. From definitions of activity coefficients and standard states, the total concentrations of organic solute and of surfactant, respectively, in either compartment of the dialysis cell, can be expressed (as shown previously32-36) as:
where [PIbt is the total solute concentration, [DDAO],ic represents the concentration of surfactant in the micelle, [DDA0lbt is the total surfactant concentration, X, is the intramicellar mole fraction of 2-phenylethanol, CPo is a limiting concentration of PEA consistent with the purecomponent standard state, CDDAO'is the concentration of monomeric DDAO in water in the absence of added PEA, (40) Berne, B. J.; Pecora, R. Dynamic light Scattering; Wiley-Interscience: New York, 1976; p 195.
Solubilization of PEA by DDAO
Langmuir, Vol. 7, No. 1, 1991 97
and ypand YDDAO are intramicellar activity coefficients of the PEA and the surfactant, based on the pure PEA and micelle standard states, respectively. The first term on the right-hand side of eqs 1 and 2 represents the concentration of monomeric surfactant (CDDAO)or unsolubilized PEA (C,), respectively. The second terms represent the concentration of these respective compounds in micellar form. If no micelles formed in the permeate in the SED experiments and if the PEA were in equilibrium across the membrane, the term ypXpCpofrom analysis of the retentate would represent tha total PEA concentration in the permeate. However, since some micelles are present in the permeate, a correction must be made, as will be discussed later. If the activity coefficients for both components in the micelle were known explicitly as functions of X,, it would be possible to solve eqs 1 and 2 simultaneously to calculate X,and [DDAOImicforboth the retentateand the permeate solutions. The activity coefficient and equilibrium constant ( K ) are related by the equation (3) The solubilization equilibrium constant is useful in calculating the amount of solute solubilized. The activity coefficient is useful because its deviation from unity indicates the favorability of the micellar environment compared to the pure PEA liquid standard state. Values of yp less than unity represent a more favorable micellar environment, and values greater than unity a less favorable environment. Hence, in this work, both yp and K will be presented. Several functional forms have been used to represent the dependence of the activity coefficient or solubilization equilibrium constant on X . We have found that the solubilization constant K often appears to vary nearly linearly with X a t values of X less than 0.2 or 0.3. A quadratic or higher order term should be included in the function if values of X as large as 0.3-0.6 are employed for typical polar solutes. A special form of this quadratic equation K = Ko(l - BX,)2
(4)
provides an excellent fit of solubilization data for phenol and for all of the mono- and dichlorophenols in aqueous hexadecylpyridinium chloride (CPC) with no significant loss of precision in fitting results as compared to using the general quadratic expression, K = Ko(l+ a X bX2),where KOis the value of the solubilization constant in the limit as X, approaches zero and B is an empirical constant. Combining eqs 3 and 4 yields
+
7,
- l/(KC,O)
= a / ( l -BX,)2
(5)
where a = l/KoCpo. If the micelle is treated as a pseudophase, the activity coefficients of the two components may be related by the Gibbs-Duhem equation. By using the Gibbs-Duhem equation and eq 5 to calculate an expression for the activity coefficient of the surfactant in the intramicellar solution, we can obtain the equation In
yDDAo
= 2 / ( 1 - B)(B In (1- X,) - In (1- BX,))
(6) to represent the dependence of the activity coefficient of DDAO in the micelle on X,.
Table I. Least-Suuares Parameters for 2-Phenylethanol (PEA) in DDAO Solution at 25 ' 0 root mean
PH
X,
Ko,M-'
B
8.5
0.0 0.25 0.5 0.75 1.0
16.8i 1.0 14.4i 1.0 12.8 i 1.0 12.1 t 1.0 17.0 f 1.0
0.757 i 0.094 0.650 i 0.079 0.696 i 0.052 0.546 i 0.059 0.783f 0.065
5.5 4.5 3.5 2.0
square deviation, 103 M 1.360 1.100 0.997
1.511 1.537
CpO (MI: 0.2226.
A nonlinear least-squares m e t h 0 d ~ l 3is~ ~utilized to analyze SED data for each system, consistent with the mathematical model described by eqs 1-6. First, eqs 1 and 2 are solved simultaneously for X , and [DDAOImic, using analytical concentration data for each of the retentate solutions, approximate trial values of the variable parameters ( B and a), known values of the standard concentrations (C,' and CDDAO'),and eqs 5 and 6 torelate the activity coefficients to the trial parameter values and X,. Next, eq 2 is employed to calculate [DDAOImi, 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. Equations 5 and 6 are used 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 experiment^.^^-^^ Finally, the predicted concentration of PEA in the permeate is inferred from eq 1for each of the experiments. The sum of squares of deviations between calculated and experimental values of [P]ht is computed in this way 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 absolute 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 to the model described above. The final column in the table lists values of the root mean square deviation (RMSD) obtained from the analysis. The individual values of K corresponding to each set of measurements were calculated from an equation derived previously,32which accounts for the presence of micelles in the permeate: K = (1- X p ) ( [ P I r e , , - [Plper,&/ ([Pl,r,,t[DDAolret,,i~ - [Plret.ht[DDAol,r,ht) (7) where the subscripts per and ret refer to the permeate and retentate compartments and the subscript tot denotes total concentration. Both [P]per,tot and [P]ret,ht are directly measurable, and [DDAO]ret,micand X, can be inferred (for the retentate solution) by solving eqs 1 and 2 simultaneously, using eq 5 (with the least-squares values of B) to relate yp and YDDAO to X,. It is assumed that the intramicellar mole fraction of PEA is the same in the permeate as in the retentate in eq 6. Finally, [DDAO]p,mi, can be calculated by using eq 2, from the measured value of [DDAO]pe,,t,t and the value of YDDAO calculated from X, and the least-squares value of B (using eq 5). Then, individual K values can be calculated by using eq 7 with results of each dialysis experiment, once the parameter B in eq 6 has been inferred from the least-squares analysis of all the data sets. (41) Christian, S. D.; Tucker, E. E. Am. Lab. 1982, 14, No.8,36. (42) Christian, S. D.; Tucker, E. E. Am. Lab. 1982, 14, No.9, 31.
98 Langmuir, Vol. 7, No. 1, 1991 E
Uchiyama et ai,
' . O 1 ' l r
PH
Figure 1. Micellar composition from the titration of 100 mM DDAO in 0.1 M NaCl with 1.0 N HCl.
Results Treating the protonated and neutral species as separate surfactants, the DDAO solution can be modeled thermodynamically as a binary mixture, the composition of which is varied by adjusting the solution pH.43 The cmc values for the mixtures of surfactants at intermediate pH levels are smaller than those for the two limiting species. A t a surfactant concentration much greater than the cmc (100 mM), almost all of the surfactant is in micellar form, so the overall composition of the surfactant mixture is virtually the same as the micelle composition. Figure 1 shows the experimentally measured overall degree of protonation obtained from titration of 100 mM solutions. The mole fraction of the cationic form increases with decreasing pH until finally only the cationic form exists a t pH 2. The presence of PEA had no measurable effect on the mole fraction of protonated surfactant in the micelle at the highest solute concentrations used in this work. From eq 4,a plot of the square root of K against X, for each system should be linear, with a slope equal to -B. Figure 2 indicates that this relationship is valid for pH values from 2.0 to 8.5. The resulting values of E are shown in Table I. Figures 3 and 4 show the activity coefficients of PEA and DDAO, respectively, plotted against X,. As X, increases, the value of y,, the activity coefficient of PEA, increases gradually toward unity. The curves must terminate a t X, values corresponding to the limiting solubility of PEA in the micelle but would reach unity at X, = 1 due to the selection of the pure liquid standard state for PEA. The activity coefficient of DDAO decreases with an increase in the mole fraction of PEA in the micelle. No consistent effect of pH on the activity coefficients of the surfactant was observed. In order to better illustrate the effect of surfactant composition on the solubilization of PEA in the micelle, the limiting solubilization equilibrium constant is shown as a function of micellar surfactant composition in Figure 5. Values of the limiting solubilization constant, KO,for the mixed surfactant systems were slightly smaller than those of the single system and showed a minimum a t a mixed molar ratio of approximately 0.5. Figure 6 shows the hydrodynamic diameter of the micelle as a function of its surfactant composition in a PEAfree system. (43) Rathman, J. F., Christian, S. D. Langmuir 1990, 6, 391.
I 0
I
I
I
I
I
0.2 0.3 0.4 MOLE FRACTION OF PEA
0.1
0.5
Figure 2. Dependence of solubilization equilibrium constants for 2-phenylethanol in dodecyldimethylamine oxide micelles on intramicellar mole fraction of solute.
O l
0
I
0.1
0.2
1
0.3
I
0.4
(
5
MOLE FRACTION OF PF!! Figure 3. Activity Coefficients of 2-phenylethanolin dodecyldimethylamine oxide micelles as a function of intramicellar mole fraction of solute.
Discussion Solubilization into Micelles Composed of Single Surfactants. It is interesting to compare the solubilization of PEA into either pure cationic or pure nonionic DDAO micelles with the solubilization of similar solutes into micelles composed of other surfactants. For the solubilization of phenol by the cationic surfactant hexadecylpyridinium chloride (CPC),32133KO = 70 M-l. For solubilization of 1-hexanol by sodium dodecyl sulfate? KO = 37 M-1; by CPC micelles, KO= 30 M-I; and into nonionic micelles composed of nonylphenol polyethoxylate, KO = 20 M-l. From Figure 2 and Table I, for cationic DDAO, KO = 17.0 M-l, and for nonionic DDAO, KO = 16.8 M-l. These DDAO solubilization constants are smaller than those for other reported alcohols and phenols in ionic or nonionic micelles. In the pure surfactant micelle, the 2-phenylethanol is thought to be solubilized with the hydroxyl group a t the surface of the micelle and the aromatic ring extending into the hydrophobic core of the micelle. In the case of the solubilization of alcohols, the hydroxyl group and the
Solubilization of PEA by DDAO
'
Langmuir, Vol. 7, No,1, 1991 99
I
0.6 0
0.1
I
I
0.2 0.3 0.4 MOLE FRACTION OF PEA
0.5
Figure 4. Dependence of the activity coefficients for dodecyldimethylamine oxide on the composition of the micelle. 25
* 41
m B
-0
0.25 0.5 0.75 1.o CATIONIC SURFACTANT/ TOTAL SURFACTANT IN MICELLE
Figure 5. Relationship between the limiting value of the sol-
ubilization equilibrium constant and the surfactant composition in the mixed micelle.
2t
01
I
I
I
J
0.25 0.5 0.75 1.0 CATIONIC SURFmANTI "TAL S U R F m A N T IN MICELLE
0
Figure 6. Micellar diameter of DDAO as a function of its surfactant composition by dynamic light scattering.
charged head groups of the ionic surfactant may be expected to attract each other by strong ion-dipole force^.^ Thus, the aforementioned alcohol solubilization constants are greater for a micelle composed of the ionic surfactant than for a nonionic surfactant micelle. In addition to this
effect, the compactness in the hydrophilic parts of the micelle may account for the fact that the solubilization equilibrium constant is less for PEA in DDAO micelles than for the other systems mentioned. The hydrophilic parts of surfactant micelles such as SDS and CPC may be less compact because of repulsion between ionic head groups due to electrostatic charge. On the other hand, since the zwitterionic surfactant DDAO has both positive and negative charges in the head group, it may form micelles in which the repulsive forces between hydrophilic head groups are weaker than in the case of ionic surfactants. The compactness of the head-group region may make it difficult for PEA to solubilize into the DDAO micelle. An alternative explanation may account for the relatively small value of KOfor PEA in DDAO micelles. Previously, it was reported that adding CH2 groups between the phenyl group and the COOH groups of carboxylic acids did not increase the tendency of these compounds to solubilize in CPC micelles.34 On the other hand, attaching aliphatic hydrocarbon groups directly to aromatic ring carbons does greatly enhance the solubilization of phenol in CPC.35144 Apparently, it is difficult for the CH2 groups in molecules like 2-phenylethanol (C6H5CH2CH20H) and hydrocinnamic acid (C6H&H~CH&OOH) to interact effectively with the hydrocarbon core of surfactant micelles, without a t the same time removing the polar groups (OH and COOH) from effective contact with the polar head groups of the surfactant. Solubilization into Mixed Micelles. The solubilization equilibrium constants for the mixed surfactant systems (DDAO/DDAOH+)are smaller than those for the individual DDAO and DDAOH+ surfactants and also smaller than would be predicted by applying linear or additive mixing rules45 to the single-component surfactant equilibrium constants. In other words, the mixed systems exhibit negative deviation from additivity. Figure 6 shows that the micelle diameter reaches a maximum for approximately an equimolar mixture of cationic and nonionic forms in the micelle. Ikeda et a1.'6 have shown that the aggregation number of the DDAO micelle reaches a maximum in the half-ionized form in the presence of 0.1 M NaC1. The compactness of the hydrophilic region of the micelle may also account for the effect of pH on the solubilizate. In the micelle containing only the cationic form, the hydrophilic region of the micelle may be less compact because the repulsive force acts between the hydrophilic groups. The same explanation may also apply for the nonionic form. The DDAO molecule possesses the dative N-0 bond; i.e., it has both positive and negative charges a t pH 8.5. Even though an attractive force acts between the positive nitrogen and the negative oxygen of adjacent surfactant molecules in the micelles of the nonionic form, repulsive forces may act between the negative atoms and/or positive atoms in the hydrophilic portion of the micelle. However, it is probably important that both the cationic and nonionic forms coexist in the mixed micelle and that these forms have equal concentrations a t a pH of approximately 5. The repulsion between the hydrophilic parts of DDAO molecules may be minimized a t pH 5 , so the mixed micelle possesses a more rigid structure than the anionic or neutral species, resulting in larger diameters and lower solubilizations. Because the values of KO of PEA in DDAO micelles (44) Smith, G. A.; Christian, S. D.; Tucker, E. E.; Scamehorn, J. F. Langmucr 1987, 3, 598. (45) Nishikido, N. J . Colloid Interface Sci. 1977, 60, 242.
100 Langmuir, Vol. 7, No. 1, 1991
depend on pH, it may be possible in commercial applications to control the volatility of this solute by changing the pH.
Acknowledgment. Financial support for this work was provided by the Department of Energy Office of Basic Energy Sciences Grant No. DE-FG05-84ER13678, De-
Uchiyama et al, partment of Energy Grant No. DE-FG01-87FE61146,
National Science Foundation Grant No. CBT-8814147, the Oklahoma Mining and Minerals Resources Research Institute, the University of Oklahoma Energy Center, Aqualon Gorp., E.I. du Pant de Nemours and CO.,KenMcGee Corp., Sandoz Chemical Co., Shell Development Co., and Unilever Corp.
