Aggregation of tetradecyltrimethylammonium bromide in water, 1,2-ethanediol, and their mixtures

Aggregation of Tetradecyltrimethylammonium Bromide in Water, 1,2-Ethanediol, and Their Mixtures SUNE B A C K L U N D , * BJORN BERGENSTAHL, t OVE M O L A N D E R , * AND T O R B J O R N W A R N H E I M ] "'1 *Department of Physical Chemistry, flbo Akademi, SF-20500 Abo, Finland, and Hnstitute for Surface Chemistry, P.O. Box 5607, S-114 86 Stockholm, Sweden Received May 25, 1988; accepted November 4, 1988 The formation of surfactant aggregates of tetradecyltrimethylammonium bromide in water, 1,2-ethanediol, and their mixtures was monitored by conductivity and density measurements. In addition, phase diagrams were determined for the systems. In water, the suffactant molality at the critical micellar concentration, CMC, is 4 ram, while, e.g., in a solvent mixture with a mass fraction of 0.40 of water and 0.60 of 1,2-ethanediol, the molality has risen to 20 ram. The phase diagrams reveal the formation of mesophases even in pure 1,2-ethanediol, showing that the surfactant aggregates in this solvent; however, no evidence for micelle formation in the solution phase was found. The decreasing sizes of the existence regions for the mesophases with increasing mass fraction of 1,2-ethanediol were found to correlate with the increasing molality at the CMC. The p h e n o m e n o n is discussed in terms of the lowered solvophobic interaction; this has been directly monitored through measuring the interfacial tension between water1,2-ethanediol mixtures and dodecane. © 1989AcademicPress,Inc.

INTRODUCTION

During the last few years, interest in studying the aggregation of surfactants in nonaqueous, polar solvents has increased considerably. Attention was drawn to the topic rather early ( 1 ), but not until now have attempts to generalize and explain the aggregation phenomena had any reasonable success. This is of course due to the current, more complete understanding of aqueous systems, but it should be emphasized that parallel research efforts can be mutually fruitful. The use of nonaqueous solvents may shed light on normal aggregation behavior in aqueous systems. The work of Evans et al. (2) on micelle formation in hydrazine has, together with studies of aqueous high-temperature micellar systems (3), challenged the conventional view of micelle formation. Micelle formation has been regarded as driven by the reduction of the "iceberg" l To w h o m correspondence should be addressed.

formation associated with the water surrounding hydrophobic solutes (4), i.e., an increase in entropy reflecting a local decrease in ordering when the hydrocarbon moieties are transferred from aqueous environment into the micelle. However, according to the work of Evans et aL (2, 3), it seems to be more or less accidential that the free energy of micelle formation in water at room temperature is dominated by the entropy term. Indeed, as pointed out by Shinoda (5), the formation of a "structured" water domain actually increases the solubility of hydrocarbon in water over what is predicted by cohesive forces alone; it is an effect that disappears at higher temperatures. Thus, the rather unique structural features of liquid water (4, 6) does not seem to be a necessary prerequisite for micelles or surfactant aggregates to occur. This has also been established in more definite terms in the study of nonaqueous lyotropic liquid crystals; e.g., lecithin forms lamellar phases in a variety of polar solvents such as 1,2-ethanediol (7),

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Journalof Colloidand InterfaceScience,Vol. 131, No. 2, September 1989

0021-9797/89 $3.00 Copyright© 1989by AcademicPress,Inc. All fightsof reproduction in any formreserved.

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BACKLUND

formamide, methylformamide (8), and even ethylammonium nitrate (9), a fused salt. While these observations certainly are of great importance, there has been rather few systematic investigations concerning the correlation of solvent properties and surfactant aggregation. One early attempt was reported by Ray (10), who with limited success tried to correlate the formation of micelles for a nonionic surfactant with a number of properties of some 20 different solvents. Neither the surface tension, dielectric constant, nor the solvent parameter (the cohesive energy density) provided a useful correlation, and Ray thus had to resort to discussions of solvent structure. Efforts to follow the change in aggregation behavior when gradually and completely exchanging the water for another polar solvent are even more rare, with some exceptions (2). The purpose of this study is to follow the change in aggregation behavior of an ionic surfactant (tetradecyltrimethylammonium bromide, TTAB) from an aqueous, micelleforming system to one where the water has been replaced by another polar solvent ( 1,2ethanediol, ED). Thus, the system water-1,2ethanediol-tetradecyltrimethylammonium bromide (H20-ED-TTAB) was investigated at six different fixed ratios of water-l,2-ethanediol (HzO-ED), from pure water to pure 1,2-ethanediol. The solution phase was characterized by different classical techniques, such as conductivity and density measurements. The characterization of the solution phase was supplemented by determinations of the phase diagrams of these systems. The occurrence of lyotropic liquid crystalline phases is an unambiguous sign of aggregated surfactant and could serve as a guideline in the investigation of the solution structure.

ET AL.

than 0.5 wt%. Freshly opened bottles of 1,2ethanediol (Aldrich, 99.5% or Riedel-de-Haen, 99%) was used as received. The water content was less than 0.2 wt%. Dodecane (Aldrich, 99%) was used as received. Water was twice distilled. Density measurements. The densities of the solutions were measured by an Anton Paar DMA 602 density meter ( 303.15 K). Conductivity. The electrical conductivities were measured by a Wayne-Kerr bridge with automatic recording of the resistances (303.2K). Phase diagram. An overview of the phase behavior was obtained by heating and mixing a number of samples ( 15-20 for each system) and equilibrating them at a few fixed temperatures in a thermostated bath. The occurrence of liquid crystalline phases was detected by viewing samples between two crossed polaroids. The more detailed phase diagrams were obtained mainly from polarization microscopy, supplemented by differential thermal analysis (DTA). Optical microscopy. Optical polarization microscopy was performed on a Reichert microscope with a fitted hot-stage, at 60× magnification. The temperature was raised 3K/ min and phase transitions were checked for reversibility. Differential thermal analysis. DTA was performed in some systems on a Mettler 2000 thermal analysis system. The temperature was raised or lowered 0.5 K/min. Phase transition temperatures were in good agreement with those found from optical microscopy. X-ray measurements. Repeat distances, d, were recorded with a slit collimated camera using Ni-filtered CuKa radiation, equipped with a position-sensitive detector Tennelee PSD-100. The equipment was repeatedly calibrated with crystalline sodium octanoate (d = 2.3 nm). The area per polar group for the MATERIALS AND METHODS surfactant in the hexagonal phase was calcu~ Chemicals. Tetradecyltrimethylammonium lated according to Mandell and Ekwall ( 11 ), bromide (Fluka, 99%) was dried overnight using the following values for the specific volunder vacuum at 320 K. The water content umes for the different components: water 1.00 in dried samples was determined to be less cm~/g, 1,2-ethanediol 0.90 cmJ/g, and tetraJournal of Colloid and Interface Science, Vol. 13i, No, 2~September 1989

395

AGGREGATION IN AQUEOUS MIXTURES 1,2

-ETHANEDIOL

method (13). The radius of the drop was obtained from photographs and the interfacial tension was calculated according to Ref. (13) (T = 293.2 K). RESULTS

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The water- 1,2-ethanediol-tetradecyltrimethylammonium bromide system forms an extensive solution phase, L, at 303 K (Fig. 1). A hexagonal liquid crystalline phase, E, occurs for mass fractions of 1,2-ethanediol in the solvent lower than 0.8 (WED < 0.8). At higher mass fractions, solid surfactant precipitates instead of the E phase. The formation of liquid crystalline phases in the same system is also shown in a series of phase diagrams at varying temperature (Figs. 2a-2f), with WEDas parameter. In pure water (Fig. 2a) a hexagonal phase, E, is observed for mass fractions of TTAB between 0.40 and 0.75 at 303 K. At lower mass fractions a solution phase L appears and at higher mass fractions and elevated temperatures a lamellar phase, D, is formed. Between the lameUar phase and

WTTAB

FIG. 1. Phase diagram of the system water-1,2-ethane-

diol ( E D ) - t e t r a d e c y l t r i m e t h y l a m m o n i u m b r o m i d e (TTAB) at 303.2 K. L denotes isotropic solution phase and E hexagonal liquid crystalline phase.

decyltrimethylammoniumbromide 0.99 cm 3/ g(12). Interfaeial tension measurements. Interfacial tension measurements between preequilibrated phases of polar solvent and dodecane were performed according to the pendent drop

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WTTAB FIG. 2. Phase diagrams of water-l,2-ethanediol, ED (at different fixed mass fractious), and tetradecyltrimethylammonium bromide, at varying temperatures. Notations as in Fig. 1, except that D denotes lamellar phase, I viscous isotropic phase, and (s) solid surfactant. Existence regions for D and E include multiphase regions with an isotropic phase. However, these are all fairly narrow. (a) Pure water as solvent; (b) WED = 0.20; (C) WED = 0.40; (d) WEn = 0.60; (e) WED = 0.80; (f) pure ED as solvent. Journal of Colloid and Interface Science, Vol. 131, No. 2, September 1989

396

BACKLUND ET AL. TABLE I

The Molality, me, at the Critical Micelle Concentration Obtained from Conductivity (Cond.) Measurements and Density (Dens.) Measurements for Solvent Mixtures with Different Mass Fractions of Water-1,2-Ethanediol

WED

me (cond.) (mmole/kg)

rn~ (dens.) (mmole/kg)

Vmon (cmS/mole)

V~ic (eroS/mole)

AV (cmS/mole)

0.0 0.2 0.4 0.6

3.83 5.8 8.8 17.9

3.9 5.2 9.0 18.0

321.3 322.6 327.8 331.9

329.9 331.5 333.4 334.1

8.6 8.9 5.6 2.2

Note. Surfactant partial molar volumes above and below me, Vmonand Vmi~,respectively, are given. The temperature is 303.2 K.

the hexagonal phase a cubic (viscous isotropic ) structure develops, i.e., an I phase. When the mass fraction of ED is increased (Figs. 2b-2f) three changes occur with respect to the existence region for the E phase: - - t h e low-temperature limit for the formation of the E phase is gradually increased - - t h e melting point decreases (directly evident only from Figs. 2c-2f). - - t h e composition region where it is stable becomes narrower.

In order to investigate the formation ofsurfactant aggregates in the solution phase, conductivity and density measurements were performed in the six different solvent systems with mass fractions of ED = 0, 0.2, 0.4, 0.6, 0.8, and 1.0, at varying molality of TTAB. The results are shown in Table I. In pure water, a distinct breakpoint in the conductivity-molality curve occurs, indicating micelle formation. The surfactant molality at the critical micelle concentration evaluated from a linear plot is 3.83 mmole/kg. This is in good agree-

Also, the phase boundary between the solid s n m2 solvated surfactant and the solution phase w gradually becomes more similar to a normal 0.57 salt-solvent phase diagram when the mass fraction of ED is increased (Figs. 2e and 2f). The geometry of the hexagonal liquid crystalline phase formed at a constant volume 0.55 fraction of surfactant equal to 0.6 was investigated by means of low-angle X-ray diffraction. The area per polar group was derived 0.53 from the repeat distance obtained ( 11 ) and is shown as a function of WEDin Fig. 3. There is :1/ I I L t I_J a significant increase, from 0.53 to 0.58 n m 2, 0.0 0.2 0.4 0.6 over the investigated interval from WED = 0.0 WED to WED = 0.55. The values for the pure water FIG. 3. Areas per polar group at a fixed volume fraction system can be compared with literature data from the water-hexadecyltrimethylammo- surfactant equal to 0.60 in the hexagonal liquid crystalline phase as a function of mass fraction of water-1,2-ethanenium bromide system; at a volume fraction of diol. T = 298 K. (©) Tetradecyltrimethylammonium surfactant equal to 0.64 the area is 0.53 bromide. (O) Hexadecyltrimethylammonium bromide n m z (14). (from (14)). Journal of ColloM and Interface Science, Vol. 131, No. 2, September 1989

AGGREGATION IN AQUEOUS MIXTURES ment with previously published values: 3.51 m m o l e / d m 3 at 303 K (15), 3.73 m m o l e / d m 3 (16), 3.79 m m o l e / k g (17), or 3.65 m m o l e / dm 3 (18) at 298 K. With increasing mass fraction of ED, the molality where the breakpoint occurs increases (Table I); in addition, it becomes less distinct. Density measurements were performed at the same mass fractions of 1,2-ethanediol as the conductivity measurements. Although less accurate from an absolute point of view for determining the CMC compared to the conductivity measurements, the breakpoints in the density curves occur at the same molalities. From the density measurements, the partial molar volume for the sufactant, Vsurf, can be determined (22). In Fig. 4, Vsurf above and below the CMC is plotted for mass fractions o f E D = 0.0, 0.2, 0.4, and 0.6, where surfactant aggregation takes place according to conducI

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397

tivity measurements. Two important features are noticeable. The difference in partial molar volume between monomeric and micellized surfactant decreases rapidly at higher mass fractions of ED than 0.2 (Table I ). The partial molar volume for the surfactant as monomer varies much more than the partial molar volume in the micellized state, suggesting, in qualitative terms, a similar environment. In addition, while Vsurfis constant except for the discontinuity at the CMC for mass fraction of ED up to 0.4, it has a gentle slope for WED = 0.6. This indicates that the pseudo-phasemodel, assuming phase separation at the CMC, fails, and that another aggregation model must be applied, in agreement with the data from the conductivity measurements. To distinguish between a proper micelle formation and a more gradual association, it has been advised to plot differential quantities vs the concentration of surfactant ( 19-21 ). In Figs. 5a and 5b the differential conductivity, AK/AmTTAB, is shown at varying molality. Although the concentration interval over which measurements have been performed in some cases is slightly too narrow, differences are revealed. For mass fractions of ED = 0.0, 0.2, and 0.4, the interval over which the differential quantity increases is comparatively narrow. Also, these curves are step-functions, suggesting that indeed a cooperative micelle formation, a precipitation of a micellar pseudophase, occurs. The solvents with WED = 0.6 and 0.8 give inflection points in the curve (note shift of scale), spread out over a wider interval, while for pure ED, the differential conductivity curve does not indicate any aggregation. Molalities are given for the critical micelle concentration of TTAB only up to WED = 0.6 (Table I) although it is, of course, possible to suggest that an inflection point occur in the curve for ED -- 0.8 Figures 5a and 5b reveal a gradual transition from a cooperative aggregation process to a less well-defined behavior ( 19-21 ). We have collected data on the solvent properties in order to correlate these with the trend of the CMCs and the existence regions of the Journal of Colloid and Interface Science, Vol. 131, No. 2, September 1989

BACKLUND ET AL.

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liquid crystalline phases. F o r ionic surfactant systems, the h y d r o p h o b i c / s o l v o p h o b i c intera c t i o n s are i m p o r t a n t t o g e t h e r with the diJournal of Colloid and Interface Science, Vol. 131, No. 2, September 1989

electric p r o p e r t i e s o f the solvent. T h e relative dielectric c o n s t a n t s for w a t e r - l , 2 - e t h a n e d i o l mixtures are available from the literature (12),

AGGREGATION IN AQUEOUS MIXTURES

while no similar data reflecting the solvophobic interaction are, to our knowledge, available. In order to obtain an estimation of that factor, the interfacial tensions between dodecane and the solvent mixtures were measured. The interfacial tension decreases in an almost linear fashion with increasing ED contents (Fig. 6). The data for the relative dielectric constant are shown in the same figure. DISCUSSION

A general point must always be kept in mind when discussing the detection of micelle formation in nonaqueous polar solvents. Claims have indeed been made that micelles form in a wide variety of polar solvents, including 1,2ethanediol (10). However, not all of the results seem consistent with what we expect to find in surfactant systems, As an example, molalities at the CMCs in the millimolal range have been reported for both sodium oleate and sodium desoxycholate in a solvent such as dimethylsulfoxide (23), which seems rather peculiar in view of their differing values in water ( 19 ). The question then becomes whether the experimental techniques used in the context really are appropriate for the purpose. Classical methods such as conductivity or surface tension measurements are only indirectly monitoring the formation of surfactant aggregates.

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399

Another important point is the distinction between micelle formation, a cooperative aggregation phenomenon, where the size distribution of the aggregates is reasonably narrow and displays a minimum, and other association processes (20). In this study we have applied different methods, normally used for aqueous systems, to detect the formation of micelles, i.e., conductivity and density measurements. By using different water-ED mixtures as solvent it is possible to follow the gradual change from the pure aqueous system, where micelles do form, to the ED system, where the solution structure is uncertain. For mass fractions of ED from 0.0 to 0.4 there is no uncertainty in claiming that a proper micelle formation takes place. Breakpoints in the conductivity and the density curves are reasonably sharp and occur at the same molalities. The phase diagram reveals a rather extensive liquid crystalline region. However, it is also clear that surfactant aggregates form even in pure 1,2-ethanediol as liquid crystalline phases (in a narrow concentration region). The molality at the CMC is increased with an increased mass fraction of ED (Table I) and the difference between the micellar state and solute state of the surfactant molecule is lowered (Table l, Fig. 4). Previous investigations of similar systems, using fluorescence quenching techniques, have shown a decrease in aggregation number when increasing the mass fraction of a polar compound other than water in the solvent. Almgren et al. (24) noted a decrease in the aggregation number from slightly below 70 in water to 55 in water-1,2ethanediol mixtures with WED ----0.20, with sodium dodecyl sulfate as surfactant, Hashimoto and Thomas (25), investigating the same system, determined an aggregation number of 29 for WED = 0.50. This information can be used to give a rough qualitative estimation of the degree of counterion association, ~. We have applied a simple model due to Evans (26), in which the aggregation number N of the micelle must be known. For TTAB in pure water, N has been estimated to be 70 (17); according Journal of Colloid and Interface Science,

Vol, 13i, No. 2~Septetnber1989

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to Evans' model this gives/3 = 0.80, in good agreement with directly measured values with ion-specific electrode (16). Assuming an aggregation number of 55 for WED = 0.20, we obtain/3 = 0.77; for a suggested aggregation number N = 35 for WEO = 0.40 we obtain/3 -- 0.65. The model is, however, not very sensitive to the aggregation number; the conductivity data show that even assuming a constant aggregation number implies a decrease in/3. We also note that the surfactant aggregates in the hexagonal liquid crystalline phase have a higher curvature and a larger area per polar group when increasing the mass fraction of ED (Fig. 3). A similar trend is observed for the hexagonal phase formed in the sodium octanoate-water-ED system (27). The tendency to form aggregates with larger areas per molecule is of course consistent with the decrease in the hydrocarbon/solvent interfacial tension (Fig. 6). We can assume that the interfacial tension at a surfactant aggregate in a solvent is related to the solvent/oil interfacial tension (although the numerical value to be used varies from model to model (28, 29)). A lower interfacial tension means a less unfavorable exposition of the surfactant alkyl chain to the polar solvent, and a larger area per polar group, other factors equal. A second factor of importance for the micellar properties, the electrostatics, tends to be slightly more confusing. A decrease in the screening of the electrostatic interactions (a reduction in the relative dielectric constant) is unfavorable to the separation of the charges of the surfactant. However, to assume this factor to be predominant is not consistent with the lowering of/3 with increasing mass fraction of ED. The effect due to the relative dielectric constant competes with the decrease in aggregate radius and an increase in head-group area, which should be coupled to a lower counterion association. The change in solvent from pure water to pure 1,2-ethanediol thus leads to several changes: --molecular solutions are favored versus micelles Journal of Colloid and Interface Science, Vol.131,No. 2, September1989

--smaller aggregates with higher curvature (and thus a more extensive hydrocarbon exposure) are favored --solution phases are favored versus hexagonal aggregates - - t h e r e is a gradual change in the phase diagram from typical Krafft-point behavior, observed in aqueous systems, to a more general salt-solvent phase diagram (30). A general explanation of the trend must, of course, be expressed in terms of the decreasing driving force for aggregation. We have discussed the correlation in terms of the solvophobic interaction, indirectly monitored through measurements of solvent/hydrocarbon interfacial tension. Correlations of similar type, such as the Gordon parameter (31 ), might also be useful in the discussion. For the appropriate theoretical modeling, for putting forth the requirements for forming surfactant aggregates, thermodynamic parameters such as free energy of transfer for the alkyl chain of the surfactant are of course required, in addition to the knowledge of the free energy balance at the micellar surface. The sometimes conflicting arguments whether surfactant aggregates form or not in a specific solvent should be tied to this point. For example, in order to obtain lyotropic mesophases in 1,2ethanediol, a surfactant alkyl chain length longer than that of the dodecyl is required (32), compared to octyl in the corresponding aqueous systems. Such considerations have curiously enough often been neglected in the literature. ACKNOWLEDGMENTS This work was financially supported in part by the Research Council at the Swedish Board for Technical Development (STUF). We thank Angela Jtnsson for performing the DTA measurements and Per Stenius for helpful comments on the manuscript. Constructive criticism from the referees is acknowledged, REFERENCES 1. Winsor, P. A., "Solvent Properties of Amphiphilic Compounds." Butterworths, London, 1954.

AGGREGATION IN AQUEOUS MIXTURES 2. (a) Ramadan, M., Evans, D. F., and Lumry, R., J. Phys. Chem. 87, 4583 ( 1983); (b) Ramadan, M., Evans, D. F., Lumry R., and Philson, S., J. Phys. Chem. 89, 3405 (1985). 3. Evans, D. F., and Wightman, P. J., J. Colloidlnterface Sci. 86, 515 (1982). 4. Franks, F., in "Water. A Comprehensive Treatise" (F. Franks, Ed.), Vol. 4, Chap. 1. Plenum, New York, 1975. 5. Shinoda, K., J. Phys. Chem. 80, 1300 (1977). 6. "Water. A Comprehensive Treatise" (F. Franks, Ed.), Vols. 1-7. Plenum, New York, 1972-1985. 7. (a) Moucharafieh, N., and Friberg, S. E., Mol. Cryst. Liq. Cryst. Lett. 49, 231 (1979); (b) E1 Nokaly, M., Ford, L. D., Friberg, S. E., and Larsen, D. W., J. Colloid Interface Sci. 84, 228 ( 1981 ); (c) Larsen, D. W., Friberg, S. E., and Christenson, H., J. Amer. Chem. Soe. 102, 6565 (1980). 8. Bergenstfihl, B., and Stenius, P., J. Phys. Chem. 91, 5944 (1987). 9. Evans, D. F., Kaler, E. W., and Benton, W. J., J. Phys. Chem. 87, 533 (1983). 10. Ray, A., Nature (London) 231, 313 ( 1971 ). 11. Maudell, L., and Ekwall, P., Acta Polytechn. Stand. Chem. 7,1, I (1968). 12. "Handbook of Chemistry and Physics," 64th ed. CRC Press, Cleveland, OH, 1982. 13. Ambvani, D. S., and Fort, T., Jr., in "Surface and Colloid Science" (R. J. Good and P. R. Stromberg, Eds.), Vol. 11, pp. 93-119. Plenum, New York, 1979. 14. Ekwall, P., Mandell, L., and Fontell K., J. Colloid Interface Sci. 29, 639 ( 1969 ). 15. Venable, R. L., and Nauman, R. V., J. Phys. Chem. 68, 3498 (1964).

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16. Zana, R., J. ColloidlnterfaceSci. 78, 330 (1980). 17. Evans, D. F., Allen, M., Ninham, B. W., and Fouda, A., J. Solution Chem. 13, 87 (1984). 18. Birdi, K. S.,Acta Chem. Scand. A 40, 319 (1986). 19. Mukerjee, P., and Mysels, K. J., "Critical Micelle Concentrations of Aqueous Surfactant Systems," Nat. Stand. Ref. Data Ser., Nat. Bur. Stand., 36, Washington DC, 1971. 20. (a) Lindman, B., and Wennerstr6m, H., in "Topics in Current Chemistry," Vol. 87. Springer-Verlag, Berlin, 1980; (b) Wennerstr6m, H., and Lindman, B., Phys. Rep. 52, 1 (1979). 21. Mukerjee, P., Mysels, K. J., and Kapuan, P., J. Phys. Chem. 71, 4166 (1967). 22. Vikholm, I., Douh6ret, G., Backlund, S., and Hoiland, H., J. Colloid Interface Sci. 116, 582 ( 1987 ). 23. Singh, H. N., Saleem, S. M., Singh, R. P., and Birdi, K. S., J. Phys. Chem. 84, 2191 (1980). 24. Almgren, M., Swarup, S., and L6froth J. E., J. Phys. Chem. 89, 4621 (1985). 25. Hashimoto, S., and Thomas, J. K., J. Amer. Chem. Soc. 105, 5230 (1983). 26. Evans, H. C., Z Chem. Soc. 579 (1956). 27. Ekwall, P., Mande11, L., and Fontell, K., J. Colloid Interface Sci. 28, 219 (1968). 28. J/3nsson, B., and Wennerstr~3m, H., Z Phys. Chem. 91, 338 (1987). 29. Eriksson, J. C., Ljunggren, S., and Henriksson U., J. Chem. Soc. Faraday Trans. 2 81, 833 (1985). 30. Laughlin, R., Adv. Liq. Cryst. 3, 42 (1978). 31. Beesley, A. H., Evans, D. F., and Laughlin, R. G., Z Phys. Chem. 92, 791 (1988). 32. W~irnheim, T., and J6nsson, A., 31 Colloid Interface Sei. 125, 627 (1988).

Journal of Colloid and Interface Science, Vol. 131, No. 2, September 1989