Langmuir 1994,10, 2170-2176
2170
Solubilization of Pentanol by Micelles of Cationic Surfactants and Binary Mixtures of Cationic Surfactants in Aqueous Solution Matthew E. Morgan,? Hirotaka Uchiyama,? Sherril D. Christian,*$? Edwin E. Tucker,? and John F. ScamehornS Institute for Applied Surfactant Research, The University of Oklahoma, Norman, Oklahoma 73019 Received December 9, 1993. In Final Form: May 2, 1994@ The solubilization of pentanol in aqueous solutions ofindividualcationic surfactants and binary mixtures of cationic surfactants at 20 "C has been investigated using head-space chromatography. Complete solubilization isotherms have been determined for pentanol in micelles of the surfactants hexadecylpyridinium chloride (CPC), trimethyltetradecylammoniumchloride (C14C1), benzyldimethyltetradecylammonium chloride (C1&zC1), benzyldimethylhexadecylammonium chloride (Cl&zCl), and hexadecyltrimethylammonium bromide (C16Cl). In every case, a large decrease in the solubilization equilibrium (or partition) constant occurs on increasing the mole fraction of pentanol in the micelles (X); a factor of 3-5 reduction in the equilibrium constant occurs as Xincreases from nearly zero to nearly 0.9. Solubilization results have also been determined for mixed micelles of C14C1and C14BzC1(system I) and CPC and ClsBzCl (system 11). Large positive synergistic effects reported previously for system I and large antisynergistic effects for system I1 are not confirmed by the results of the present-studies. For both of these mixed micellar systems, small decreases in the value of the solubilization equilibrium constant, relative to values predicted with an assumed additivity relationship, are observed.
Introduction There have been many studies of the extent of solubilization of aliphatic alcohols by individual surfactant micelles in aqueous solution and several investigations of the solubilization of these compounds by binary mixtures of surfactants of different types (anionic plus nonionic, cationic plus nonionic).'-13 Recently, data have been reported for the solubilization of pentanol in mixed micelles of cationic surfactants, with the unusual result that the extent of solubilization (for a given total concentration of surfactants and penanol) in some cases apparently exhibits large departures from additivity.11J2 In the case ofmixed micelles of C1&1 and C14BzC1 (system I), large enhancements in the solubilization of pentanol (a positive synergistic effect) were found a t mole fractions of C1azC1 in the mixed micelle of approximately 0.1." Very recently, it was reported that quite large negative deviations from additivity occur throughout wide concentration ranges
* To whom correspondence m a y be addressed. Department of Chemistry and Biochemistry, The University of Oklahoma, Norman, OK 73019. School of Chemical Engineering and Materials Science, The University of Oklahoma, Norman, OK 73019. Abstract published in Advance A C S Abstracts, June 15,1994. (1)Hayase, K.; Hayano, S. Bull. Chem. SOC.Jpn. 1977,50, 83. (2) De Lisi, R.; Genova, C.; Turco, V. J.Colloid Znteface Sci. 1983, 95,428. (3) Treiner, C.;Khodja, A.; Fromon, M.; Chevalet,J . J . Solution Chem. +
*
@
1989,18, 217.
(4) Bury, R.; Treiner, C. J. Solution Chem. 1989,18, 499. (5)Nguyen, D.; Vanable, R.; Bertrand, G. Colloid S u f . 1992,65, 231. (6) Rao, I. V.; Ruckenstein, E. J. Colloid Znterface Sci. 1986,113, 375. (7) Treiner, C.; Bocquet, J.; Pommier, C. J. Phys. Chem. 1986,90, 3052. (8)Treiner, C.; Khodja, A.; Fromon, M. Langmuir 1987,3,729. (9) Bury, R.; Treiner, C. J.Phys. Chem. 1991,95, 3824. (10)Njpyen, C. M.; Scamehorn, J. F.; Christian, S. D. Colloid Surf. 1988,30,335. (11)Bury, R.; Treiner, C.; Chevalet, J.; Makaysi, A. Anal. Chim. Acta 1991,251,69. (12) Treiner, C. 204th National Meeting of the American Chemical Society, Washington, DC August 1992; Abstract COLL128. (13)Nishikido, N.; Sugihara, G. In Mixed Surfactant Systems; Ogino, K, Abe, M., Eds.; Marcel Dekker, New York, 1993; p 393.
for pentanol in mixed micelles of CPC and C&zCl (system II).12 One implication of the large (positive) synergistic effects in solubilization, should they occur in mixtures of surfactants of the same or different charge types, is that it may be possible to use carefully chosen mixtures of surfactants in micellar separation methods to achieve specificity in removing particular solutes from aqueous process streams or other contaminated water. Therefore it will be important to determine if significant cooperative solubilization effects actually do occur in certain types of systems and, if such effects are confirmed, to develop methods for predicting when pronounced synergism will occur. One area of research (in the case of single surfactant systems) that seems to have received relatively little attention is the determination of the effect that intramicellar composition (mole fraction of aliphatic alcohol in the micelle orX) can have on the solubilization equilibrium constant (or partition ratio) for the alcohol. An important goal of many solubilization studies from our laboratories has been to determine complete "solubilization isot h e r m ~ (K " ~as ~ a function of X , from X 0 to values of X near saturation). Such data are valuable in probing the nature of molecular interactions between the micellar surfactant and organic solubilizate. They can also be used to predict the effectiveness of micellar-based separation methods, including micellar-enhanced ~ltrafiltrationl~ (MEW), which can be used to remove organic contaminants from aqueous streams. In the case of polar solutes such as phenol and substituted phenols, benzaldehydes, etc., bound by cationic micelle^,'^-^^ solubilization isotherms are quite accurately
-
(14)Nguyen, C. M.; Christian, S. D.; Scamehom, J. F. Tenside, Surfactants, Deterg. 1988,25,6. (15)Christian, S. D.; Scamehom, J. F. In SuTfactant-Based Separation Processes; Scamehorn, J. F., Harwell, J. H., Eds.; Marcel Dekker, New York, 1989; p 3. (16) Lee, B.-H.; Christian, S. D.; Tucker, E. E.; Scamehon, J. F. Langmuir 1990,6, 230. (17) Lee, B.-H.; Christian, S. D.; Tucker, E. E.; Scamehorn, J. F. J. Phys. Chem. 1991,95,360.
0743-746319412410-2170$04.50/0 0 1994 American Chemical Society
Solubilization of Pentanol by Cationic Surfactants
Langmuir, Vol. 10,No. 7, 1994 2171
represented by the equation K = Ko(l - B W , where KO represents the limiting solubilizationequilibrium constant for the solute as X approaches 0.16J7~20~21 The parameter B is related to the number of head groups in the micelle that constitutes a binding site, to which the solute molecules presumably attach by Langmuir-type adsorption. We thought it would be of interest to determine if solubilization data for pentanol in the individual cationic surfactant micelles could also be correlated reliably by this simple equation.
h
*
v
1c
t
A
water 0
CPC 0.05M
Experimental Section Trimethyltetradecylammoniumchloride (C14Cl)was obtained from TCI, Japan; 98% purity was ascertained using fast atom bombardment mass spectrometry. Benzyldimethyltetradecylammonium chloride (C14BzCl)and benzyldimethylhexadecylpyridinium chloride (C16BzCl)were obtained from Sigma, and were 99% pure and 97% pure, respectively. Hexadecylpyridinium chloride, CPC, was obtained from Hexcel Co.; 99% purity was deduced from the lack of a surface tension minimum in plots of surface tension versus concentration. Hexadecyltrimethylammonium bromide (CTAB)was recrystallized from ethanol several times. 1-Pentanol, obtained from Sigma, was greater than 99% pure. All solutions were prepared with doubly distilled and deionized water. Head-space gas chromatography (HSGC)was used to measure the partial pressure of 1-pentanol in the vapor in equilibrium with aqueous solutions of known composition. A Perkin-Elmer Model Sigma 300 gas chromatograph with a flame ionization detector was used for these measurements. The column used was a Tenax GC, with a packing material made of a porous polymer based on 2,6-diphenyl-p-phenylene oxide (columnlength 3 ft., width 0.d. l/s in., mesh range SO/lOO). Peak areas were recorded on a Varian 4270 integrator. Surfactantlpentanol solutions were prepared, and 20-mL samples of solution were placed in EPA certified 40-mL sample vials, sealed airtight with silicon septa, and allowed to equilibrate in a water bath at 20 "C for 24 h before injection. The gas chromatograph, oven, injector, detector, and column temperatures were at 200 "C. A gastight syringe was used, and all samples were injected manually. Approximately 1mL of sample was drawn out of the sample vial, the sample volume was adjusted to 0.5 mL, and then the sample was injected into the gas chromatograph. We have been able to analyze the vapor pressure results to infer the partitioning of pentanol between vapor and liquid phases and to determine the concentrations of pentanol in the "bulk" aqueous solution and of pentanol solubilized by surfactant micelles. The partial pressure of pentanol in the vapor phase was obtained from the activity of pentanol in the vapor phase, defined as the ratio of concentration of the solute in the vapor phase above the solution to that above pure liquid pentanol. The vapor pressure of pure pentanol at 20 "C used in this calculation is 1.69 Torr.24
Results Preliminary to performing the solubilization study, we carried out a series of "blank" experiments to obtain the dependence of the partial pressure of pentanol on its mole fraction or molarity in the aqueous solution phase. Information about the partial pressure of pentanol, which varies nearly linearly with the molar composition of
0
0.05
0.1
0.15
0.2
0.25
0.3
[Pentanol] (M) Figure 1. Partial pressure ofpentanol as a function of molarity for a solution in water and a cetylpyridinium chloride 0.05 M solution at 20 "C.
pentanol in the very dilute solutions studied (Le.,nearly obeys Henry's law), is of course required in interpreting HSGC data for pentanol in the surfactant solutions. Figure 1 shows plots of the partial pressure of pentanol as a function of the concentration of pentanol in water and in 0.05 M CPC at 20 "C. The partial pressure of pentanol increases with the increasing concentration of pentanol. A second-order equation was fitted to the data (vapor pressure vs pentanol concentration in water), and this function was used to calculate the concentration of unsolubilized pentanol in solution from a given measured pentanol activity (inferredfrom measured partial pressure values). U P = 3.270Cp,b,k -
1.368CP,,d2
The variation of the partial pressure of pentanol for pentanoYwater solutions with concentration indicates that the activity of pentanol does not vary exactly linearly with the pentanol concentration. The deviation from linear Henry's law behavior can be attributed to the selfassociation ofpentanol to form dimers and perhaps higher molecular weight aggregates. The actual activity change can be quantified by the Hansen-Miller equation^,^^^^^ which can be used to relate the activity coefficients for a solute to the mole fraction of solute in solution and empirical constants. At any given partial pressure of pentanol (Le., at constant pentanol activity), the difference between the concentration of pentanol in a surfactant solution and the concentration of pentanol in pure water at the same activity (neglecting any salting-out effect)corresponds to the concentration of pentanol solubilized in the micelle. It is, therefore, possible to calculate the solubilization equilibrium constant, which expresses the partitioning of pentanol between micellar and bulk phases, almost directly from the HSGC data. We define the solubilization equilibrium constant by
(18)Lee, B.-H.; Christian, S. D.; Tucker, E. E.; Scamehorn, J. F. Langmuir 1991,7 , 1332.
(19)Uchiyama,H.;Christian,S.D.;Scamehorn,J.F.;Abe,M.;Ogino, K. Langmuir 1990,6,95. (20)Kondo, Y.;Abe, M.; Ogino,IC;Uchiyama, H.; Christian, S. D.; Tucker, E.E.;Scamehorn, J. F. Langmuir 1993,9,899. (21)Christian, S.D.: Tucker, E. E.: Scamehorn, J. F.: Uchiyama, H.
Colloid Polym. Sci., in press. (22)Uchiyama, H.; Christian, S. D.; Tucker, E. E.; Scamehom, J. F. J.Phys. Chem., in press. (23)Uchiyama, H.; Christian, S. D.; Tucker, E. E.; Scamehorn, J. F. J . Phys. Chem., submitted for publication. (24)Butler, J.; Ramchandani, C.; Thomson, D. J.Chem. SOC.1935, 138, 280.
is the intramicellar mole fraction of pentanol, Cp,mon is the molarity of monomeric pentanol in the bulk aqueous phase (25)Corrin, M.; Harkins, W. J.Am. Chem. SOC.1947,69,683. (26)Tipton, R.M.S.Thesis, The University of Oklahoma, 1989. (27)Morgan, M.E.M.S. Thesis, The University of Oklahoma, 1993.
2172 Langmuir,Vol. 10, No. 7,1994
Morgan et al.
Table 1. Values of the Critical Micelle Concentration,Corrin-Harlrins Parameters,Limiting Solubilization Equilibrium Constant, and Empirical Parameter B Obtained from Eauation 8 surfactant CMC, M B U. KO,M-' B 2B CPC 9.0 10-4 16 -0.73% -5.286 20.4 f 0.8 0.969f 0.051 1.94 ~~
ClSBZCl
4.0 x 10-436 9.2 1 0 - 4 3 7 5.4 10-3 36 2.2 10-3 36
CTAB
c14c1 C14BzCl
-0.73% -0.73% -0.70% -0.7038
16.6 f 0.6 14.8f 1.2 14.9f 0.7 8.1 f 0.6
-5.878 -5.253 -3.862 -4.535
0.728f 0.025 0.815 f 0.1 0.648f 0.038 0.48 f 0.062
1.46 1.63 1.30 0.96
25 0
K, Lo CTAB 0.0s M
20
15 10
5
"0
0.2
0.4
0.6
"0
0.2
0.4
XP 1.0
,. # a .Y
t
.
0.8
1.0
- 7,=l/K,C; 0
- y,=l/K,C,'
7, inCPC in CPC micelk
0 c)
e .Y
6
0.6
s
7, la CTAB micelle 0
0.8 0.6
8
h
h
.E 0.4
.S
.e
.?; c)
z
0.8
0.6
XP
0.4
c)
8
0.2 0' 0
0.2
0.4
0.6
0
J
t
0
XP
0.2
0.4
0.6
0.8
XP
Figure 2. (a, top) Solubilizationconstant as a function of the mole fraction of pentanol in CPC micelles. (b,bottom)Activity coefficient as a function of the mole fraction of pentanol in CPC
micelles.
(inferred from HSGC data for the aqueous solution of the pentanol by eq l), Cp,micis the concentration ofthe pentanol in the micellar form, and Csd,bt and Cmon denote the concentrations of the surfactant in the system and in the monomer form, respectively. The effective monomeric concentration, or activity, of an ionic surfactant is lowered both by the addition of organic solutes and by an increase in concentration of counterions in solution. In this study, in the absence of added organic solute, the monomer concentrations of surfactant are inferred from the CorrinHarkins equationz5 log Cmon = -B log Ccl,fiW - a
0.2
(4)
where Ccl,freeis the concentration of free counterion (C1-) and a and j3 are empirical constants which can be obtained by measuring the critical micelle concentration (CMC)of a surfactant in the presence of different concentrations of electrolyte.26 The Corrin-Harkins parameters used in this study are listed in Table 1. Given the value of , C it is possible to utilize information about the intramicellar composition and estimates of the activity coefficient of pentanol in the micelle to calculate the value of Cmo,for each solution containing both pentanol and surfactant. The iterative procedure used to obtain the corrected C ,,, values is described in refs 16, 17, 19, and 27.
Figure 3. (a, top) Solubilization constant as a function of the mole fraction of pentanol in CTAB micelles. (b,bottom)Activity coefficientplotted as a function of the mole fraction of pentanol in CTAB micelles.
Figures 2a, 3a, 4a, 5a, and 6a show the dependence of the solubilization equilibrium constants on the mole fraction of pentanol in CPC, CTAB, CIeBzCl, C14Cl, and C1J3zC1 micelles, respectively. The solubilization constants decrease considerably and monotonically with the increasing mole fraction of pentanol in the micelle. The activity coefficient of pentanol ( y ) is equal to the ratio of the fugacity of the pentanol in the micellar solution ( P ) to the fugacity of pure pentanol (Po)at the given temperature, divided by X. (The fugacity and partial pressure of pentanol are assumed to be equal.) Figures 2b, 3b, 4b, 5b, and 6b show the relationship between the activity coefficient of pentanol and the mole fraction of pentanol in the micelle. The activity coefficients of pentanol in the surfactant micelle increase toward unity with the increasing mole fraction of pentanol in the micelle. In a binary surfactant system, the two surfactants form mixed micelles, in which the compositioncan be somewhat different from the mixing ratio in the solution. In calculating the solubilization equilibrium constant of pentanol in a binary surfactant mixture, one should consider the amount of the surfactants in the mixed micelle, as well as in the monomeric form, to obtain an accurate mole fraction of pentanol in the micelle.
Solubilization of Pentanol by Cationic Surfactants
Langmuir, Vol. 10, No. 7,1994 2173 25 1
20
01 0
" " " " " "
0.1
0.2
" ' " " ' " ' ' ' ~
0.3
0.4
0.5
O
0.6
I
- K,=K,(I.Bx,)*
1
0
0
t
"
" " 0.2
1.0
1
1.0
e
o'""'~"""'"''"'""'' 0.2
0.3
0.8
0.6
0.8
r
.
0.1
0.6
XP
XP
0
0.4
K, in CI,CI 0.05 M
0.4
0.5
0.8 -
0
0.6
XP Figure 4. (a, top) Solubilization constant as a function of the mole fraction of pentanol in C16BzC1 micelles. (b, bottom) Activity coefficient plotted as a function of the mole fraction of pentanol in C16BzC1 micelles.
The mole fraction of pentanol in the mixed micelle is given by (5) where
Cmic1and Cmic2are the concentrations of surfactants 1and 2 in the micellar form, Csurfl,tot and CsurfZ,tot are the total concentrations of surfactants 1and 2 in the solution, and Cmonl and Cmon2 denote the monomer concentrations of surfactants 1 and 2, respectively. In this analysis, the concentration of each monomeric surfactant in equilibrium with the mixed micelle can be inferred by multiplying the CMC values calculated for the individual surfactants, which are consistent with the Corrin-Harkins by the mole fraction of each surfactant in the micelle and by the activity coefficients calculated from the Rubingh applied to mixed micelles using an interaction parameter (&). The /3m values for the two mixed surfactant systems studied here are -0.50 for C1&1 C14BzC1 and -0.8 for CPC + c16BzC1.ll All equations are solved using a simultaneous equation solver. However, the differences between the solubilization equilibrium constants for surfactant mixtures, calculated with and without corrected monomeric surfactant concentrations, were so small that they are the same within experimental er1-01.~'
+
(28) Rubingh, D. In Solution Chemistry of Surfactants; Mittal, K., Ed.; Plenum Press, New York, 1979;p 37.
0
- y,=l/K,C; 0
y, In C,&I micelle
0.2
0.4
XP
Figure 6. (a, top) Solubilization constant as a function of the mole fraction of pentanol in c14c1micelles. (b,bottom) Activity coefficient plotted as a function of the mole fraction of pentanol in C14C1 micelles.
Figure 7 depicts the concentration of pentanol not solubilized in the mixed micelle as a function of the solution mole fraction of C1&zC1 in CPC/C16BzC1 mixed systems. The concentration of monomeric pentanol in the bulk phase is observed to be almost constant throughout the total mole fraction range for C16BzCl and CPC. Using the concentrationof unsolubilized pentanol, the solubilization equilibrium constants of pentanol in the mixed micelle are calculated. Figure 8 shows the solubilization equilibrium constants in CPC/C~GBZC~ mixed surfactant systems as a function of the solution mole fraction O f C16BzCl in the mixed solution. The solubilization constants decrease slightly with the increasing mole fraction of C16BzCl at low relative concentrations of that surfactant. Figure 9 is a plot of the unsolubilized pentanol concentration in the C14CVC14BzCl mixed surfactant system as a function of the solution mole fraction of C14BzC1. The concentration of pentanol in the bulk phase increases slightly with an increase in the mole fraction of Cl4BzC1 in the solution. Values of the solubilization equilibrium constant of pentanol in the c&l/C~BzCl mixed surfactant system are shown in Figure 10. The solubilization constant decreases slightly with the increasing mole fraction of CJ3zC1 in the solution. Also, using the limiting solubilization equilibrium constant (KO) and B parameter for pentanol in micelles of pure C&l, the solubilization constant is calculated at the pentanol concentration of 0.05 M and shown in the figure. Discussion
Equilibrium solubilization results are intrinsically valuable, providing fundamental information about the
Morgan et al.
2174 Langmuir, Vol. 10,No.7,1994 25 1
- K,=K,(I-Bx,)' 0
20 -
K, In CI,BzCI0.05 M
c)
c (D
c) v1
e
8 e
.5!
0 0
15 -
8
e .5!
10 -
.5 .P
5 -
CPC/C,,BzCI total conc. 0.05 M PentanolO.05 M
"0
0.2
I
3
I
I 0
0.2
0.6
0.4
0 m
0.8
1.0
0.8
Figure 8. Solubilization constant as a function of the solution mole fraction of CI&zCl in CPC/Cl&zCl mixed surfactant solutions (total surfactant concentration 0.05 M, pentanol concentration 0.05 M).
/ I
1.0 1
0.8
0.05 1
0
b 0.6
0.6
XC,,BzCISO1
XP
c .P .Y
0.4
1
I)
Q)
0 0
h .Z
.L
0.4
c)
3
0.2
0
a
D
0
0.6
0.4
0.2
0.8
XP Figure 6. (a, top) Solubilization constant as a function of the mole fraction of pentanol in C1J3zC1 micelles. (b, bottom) Activity coefficient plotted as a function of the mole fraction of pentanol in C1J3zC1 micelles. 0.05
1
1
.P
I
F 0.02
C,,CUC,,BzCI total conc. 0.05 M PentanolO.05 M
s:c 0.01
& 0 0.4 0.8 0
0.2
0.6
1.0
XC,,BZCISOi
Figure 9. Unsolubilized pentanol monomer concentrationas a function of the solution mole fraction of C14BzC1 in C&l/ C1J3zC1 mixed surfactant solutions (total surfactant concentration 0.05 M, pentanol concentration 0.05 M). 15 I
I
Estimated Ks uslng Eq 8
1 0
0.2
0.4
0.6
0.8
C14CUC14BzCI total conc. 0.05 M PentanolO.05 M
1.0
xc1,BzcII0~ Figure 7. Unsolubilized pentanol monomer concentration as a function of the solution mole fraction of cl&zcl in cPc/c16BzCl mixed surfactant solutions(total surfactant concentration 0.05 M, pentanol concentration 0.05 M).
01
0
"
0.2
"
0.4
"
0.6
"
0.8
'
I
1.0
XC,,BzCIIOi
F'igure 10. Solubilizationconstant as a function of the solution mole fraction of C18zC1 in C&l/C1&Cl mixed surfactant solutions (total surfactant concentration 0.05 M, pentanol concentration 0.05 M).
interactions of organic solutes with surfactant micelles and the locus of solubilization of molecules within micelles. Polar organic solutes such as the aliphatic alcohols, carboxylic acids, phenols, and cresols tend to solubilize polar molecules are anchored in the surface region of with a strong interaction between their polar groups and the polar/ionic head groups of s u r f a c t a n t ~ . ' ~ J Solu~ ~ ~ ~ - ~ ~surfactant micelles and that their aliphatic or aromatic moieties extend at least partly into the hydrocarbon core. bilization results are generally consistent with other Typically, when solubilized in the common ionic and physical evidence, indicating that the head groups ofthese nonionic micelles, these solute species have activity coefficient values on the order of 0.1 or less (based upon (29)Bhat, S. N.; Smith, G. A.; Tucker, E. E.; Christian, S. D.; Scamehom, J. F. 2nd. Eng. Chem. Res. 1987,26,1217. pure component standard states), and these values tend (30) Smith, G. A.; Christian, 5. D.; Tucker, E. E.; Scamehom, J. F. gradually to increase toward unity asX(the mole fraction J . Solution Chem. 1986,15, 519. of solubilizate in the micelle) i n ~ r e a s e s . l ~As - ~required ~ (31) Higazy, W. S.; Mahmoud, F. Z.; Taha, A. A.; Christian, S. D. J . by the reciprocal relationship between the solubilization Solution Chem. 1987,17, 191.
Langmuir, Vol. 10, No. 7, 1994 2175
Solubilization of Pentanol by Cationic Surfactants equilibrium constant and activity coeficient, Kvalues for these solutes decrease substantially as X increases. In marked contrast, highly hydrophobic solutes such as the aliphatic hydrocarbons solubilize primarily within the hydrocarbon core region of typical aqueous micelles. Activity coefficientsfor these species tend to decrease from values in the range 3-10 at infinite dilution to values nearer 1 as X increases.16 The aromatic hydrocarbons are intermediate in behavior between highly polar solutes, which are anchored in the micelle surface region, and the aliphatic hydrocarbons. Thus, solutes such as benzene and toluene have activity coefficients close to unity in a variety of surfactant micelles, perhaps reflecting their tendency to solubilize both in the head-group and hydrocarbon core regions of typical micelles. Moreover, the excess enthalpy and entropy changes for benzene, transferring from the pure solvent phase into typical ionic surfactant micelles (for pure component standard states), are small and usually p o s i t i ~ e . ~ ~ , ~ ~ The results reported here (Figures 1-6) for pentanol in CPC, CTAB, C16BzC1, C&1, and C18zC1micelles are quite analogous to solubilization data for phenol or phenol derivatives in ionic surfactant micelles. Values of the activity coefficient for pentanol are smaller than unity, approaching unity with the increasing mole fraction of pentanol in the micelles; correspondingly, the K values (which vary reciprocally with values of the activity coefficient)decrease with increasing X,. The fact that K decreases significantly as X increases suggests that a Langmuir-type adsorption model may be useful in correlating the solubilization results.34 Thus, it is argued that the observed decrease in Kwith increasingximplies that the binding of pentanol -OH groups at the surface of the micelle uses up some of the adsorptive capacity of the cationic surfactant micelles. We have fitted the K vs X data for each system to the semiempirical equation
K = Ko(l - BXJ2= Ko(l- 2BX,
+ PX;)
(8)
proposed by Lee et aZ.15-23The equation is consistent with the Langmuir adsorption e q u a t i ~ nat~ small ~ , ~ ~values of X, assigning a certain number of surfactant molecules in the micelle to a single adsorption site. The negative of the coefficient of theX term (that is, twice the value of the parameter B) in eq 8 can be shown to be equal t o the number of surfactant molecules constituting one "site" in the Langmuir adsorption mode1.16J7 The positive value of the quadratic term (B2XP2)causes the calculated solubilization isotherm to decreases less rapidly than is predicted by the limiting linear form of eq 8, K = Ko(1 2BXp),which would require that K approach zero as X, approaches 142B). The values of KO and B in Table 1have been inferred by fitting all of the solubilization data to eq 8. Thus, the numbers of surfactant molecules constituting one site (2B) vary from 1.6 to 1.9 for the c16 surfactants, decreasing to 1.2-1.0 for the C14 alkyl chain surfactants. Previously, we compared the binding of phenols and chlorinated phenols to CPC micelles, for which 2B is on the order of 2-3, with that of solutes like the benzaldehydes,for which 2B is on the order of 1-1.5.16J7 The smaller values of 2B for the benzaldehydes, as compared with the phenols, were interpreted as indicating that the benzaldehydes are less strongly bound in the head-group region of CPC than are (32) Christian, S.D.; Tucker, E. E.; Smith, G. A.; Bushong, D. S.J . Colloid Interface Sci. 1986, 113, 439. (33) Smith, G.A.;Christian, S.D.; Tucker, E. E.; Scamehom, J. F. J . Colloid Interface Sci. 1989, 130, 254. (34) Mukejee, P.J . Pharm. Sci. 1971,60,1531.
the more polar amphiphiles. Values of 2B in the range 1-2 for the pentanollsurfactant systems may imply that pentanol also interacts less strongly with the pyridinium or quaternary ammonium groups near the micellar surface than do phenol and the chlorinated phenols. alkyl chain surfactants, CPC had the Of the three largest KOand 2B values. The moderate differences in solubilization at low mole fractions can be attributed to the different head groups on each surfactant. The pyridinium cation in the CPC head group is apparently more effective in solubilizingpentanol at low mole fractions than is the benzyl group, presumably because of a strong ion-dipole interaction between the pyridinium group and the hydroxy group of the pentanol. On the other hand, the aromatic group in the C16BzCl molecule may sterically oppose penetration of the pentanol molecule into the micelle. CTAB, having no aromatic groups in its hydrophilic moiety, is intermediate in behavior between CPC and C&zC1 at low mole fractions of pentanol. The relative values of K for the C16BzC1 and C14BzC1 systems indicate that surfactants with longer hydrophobic chains consistently solubilize more pentanol than the shorter-tailed surfactants. This result is consistent with other solubilization results for homologous series of surfactants, such as sodium alkyl sulfates.35 In recent papers by Treiner et aZ.,11J2calorimetric results led to the conclusion that the extent of solubilization of pentanol exhibits a sharp maximum in the vicinity of 0.1 mole fraction of C14BzC1 in the mixtures of C14C1 and c14BzC1. Thus, from results in their paper, the solubilization equilibrium constant can be calculated to be approximately 48 M-' at the maximum (mole fraction 0.1 of C14BzCl). In the present study, in order to check this conclusion, we prepared a large number of solutions in this mole fraction region, so as not to overlook any indication of synergistic solubilization. The reported synergism is not reproduced by the present study; that is, there is no sharp increase in solubilization at 0.1 mole fraction, or in any other mole fraction region. There has been discussion in the literature2 about discrepancies between solubilization results determined using different experimentaltechniques,for example headspace chromatography and the calorimetric method used by Treiner et al. and other investigators.2J1J2The different methods involve quite different experimental procedures and assumptions in the analysis of data, although in our opinion the head-space chromatography method (and vapor pressure methods in general) requires fewer assumptions in leading from the observed data to calculated solubilizationisotherms than the calorimetricmethods."J2 However we are unable to explain why calorimetry experiments should have yielded such abrupt changes in the partition ratio (or solubilization equilibrium constant) in a narrow range of relative concentrations of the two cationic surfactants. The head-space chromatography results obtained in the present study for C14C1 Cl4BzC1are very similar to HSGC data reported previously for hexanol solubilizedin similar mixed micelles.1° The value of K,is initially about 11.5 M-' for pure C14C1, and decreases with the addition of C14BzC1 to a value of about 8.5 M-' at a solution mole fraction of 0.1 C14BzC1. However, from the single surfactant data for C&l, K,was calculated to be closer to 10
+
(35) Rosen, M. J. Surfactants and Interfacial Phenomena, 2nd ed.; John Wiley & Sons: New York, 1988; p 176. (36) Treiner, C.; Makayssi, A. Langmuir 1992, 8, 794. (37) Rosen, M. J. Surfactants and Interfacial Phenomena, 2nd ed.; John Wiley & Sons: New York, 1988; p 126. (38)Bushong, D. S.Ph.D. Dissertation,The University of Oklahoma, 1985.
Morgan et al.
2176 Langmuir, VoE. 10, No. 7,1994
M-l than 11.5M-l, which implies that the initial decrease of K, is more gradual than would be indicated from the mixed surfactant data alone. The value of K, remains nearly constant at Xc1fiZc1 values from 0.2 to 1.0. The initial decrease in K, might be explained by a competition between pentanol and C14BzC1 for solubilization sites in the palisade layer of the predominantly C14Cl micelle. As the mole fraction of C1J3zCl increases, the micelle takes on the solubilizationcharacteristics of a micelle composed of pure C14BzC1. The new solubilization data are inherently interesting, but they do not verify the report of a pronounced synergistic solubilization effect. A strong antisynergistic solubilization effect was reported previously for the mixture of CPC Cl&zC1.12 The solubilizationof pentanol in this mixture was observed to decrease dramaticallyin the solutionmole fraction range ofO.l-0.9 C16BzCl. Figure 8from the present study shows a much more gradual change in K, with a variation in solution mole fraction, and the values of K, are all intermediate between the K,values for individual surfactants. As in the case of C14C1 C1J3zC1, there is no observed evidence for either pronounced synergistic or antisynergistic effects. An accurate mathematical model is important in inferring solute solubilization constants or partition coefficients. All si&icant factors that affect the experimental results need to be included in the model. Changes in surfactant monomer and micellar concentrations, counterion effects, and pentanol self-association cannot be ignored in processing experimental data to derive meaningful information. Our knowledge has progressed beyond the point where we need only subtract the CMC from the total surfactant concentration to get the concentration of surfactant in micelles. Such a simplification
+
+
may produce inaccurate solubilizationresults and lead to erroneous conclusions. It should be emphasized that each of the corrections applied to infer the surfactant monomer concentration, Cmon (for counterion concentration, surfactant concentration, and presence of the organic solubilizate), decreases Cmon. Cmon becomes SO small in some cases, for example in the 0.05 M c16 surfactant solutions, that it is a reasonably good approximation to ignore the presence of Cmonin the solubilization calculations. In fact, when the K, values are calculated for the mixed surfactant systems, both by using the corrected Cmonvalues and with the assumption that Cmon= 0, the discrepancy between the two sets of values is well within experimental error. On the other hand, for the 14-carbon single component surfactants this difference does become significant, particularly at low molarities of the ~ u r f a c t a n t . ~ ~
Acknowledgment. The authors 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, National Science Foundation Grant CBT 8814147, University of Nevada-Reno, Contract 90-09, an Applied Research Grant from the Oklahoma Centers for the Advancement of Science and Technology, and the Center for Waste Reduction Technologies of the American Institute of Chemical Engineers, Agreement No. N12-Nl0. In addition, they gratefully acknowledge the assistance of industrial sponsors of the Institute for Applied Surfactant Research, including Aqualon Co., Kerr-McGee Corp., Sandoz Chemicals Corp., E. I. duPont de Nemours & Co., Unilever, Inc., Union Carbide Corp., Dow Chemical Co., IC1 Chemicals, and Shell Development Co.
