Binding of Aromatic Anions to Cetylpyridinium Aggregates Either

4-aminosalicylate, cetylpyridinium 5-aminosalicylate, and cetylpyridinium chloride has been investigated at silica/water interfaces using adsorption i...
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Langmuir 1998, 14, 7493-7502

7493

Binding of Aromatic Anions to Cetylpyridinium Aggregates Either Adsorbed at Silica/Water, Alumina/Water, Titanium Dioxide/Water Interfaces or in Solution P. Favoriti and C. Treiner* Laboratoire Liquides Ioniques et Interfaces Charge´ es, UMR CNRS 7612, Universite´ Pierre et Marie Curie, 4, Place Jussieu, Bat.74, BP 51 Paris 75005, France Received July 14, 1998. In Final Form: October 6, 1998 The influence of the counterions on the adsorption properties of cetylpyridinium salicylate, cetylpyridinium 4-aminosalicylate, cetylpyridinium 5-aminosalicylate, and cetylpyridinium chloride has been investigated at silica/water interfaces using adsorption isotherm determinations and electrophoretic and surface tension measurements. Earlier results are confirmed which showed that replacing the chloride ion by the salicylate ion results in a 6-fold increase in the plateau value of the surfactant ion. It is suggested that this is the consequence of a close packing of the cationic surfactants when associated to the salicylate ion at either the solid/water or air/water interface. The results are compatible with the formation of a surfactant monolayer with the headgroups facing the solid surface. The binding constants of the organic anions to the cationic headgroups have been evaluated using a Langmuir type isotherm and were shown to be larger for adsorbed aggregates as compared to free micelles by a factor of about 3. Although the binding constant is larger by a factor of 10 for the salicylate ion as compared to the aminosalicylate derivatives, the surfactant adsorption isotherms and the electrophoretic mobilities are similar for the three cetylpyridinium salts. It is suggested that the unusually large adsorption of these surfactants at the silica/water interface is related to the surface stacking of the aromatic counterions which favors the close packing of the surfactant ions. Results obtained at the titanium dioxide/water and the alumina/water interfaces do not contradict these conclusions. The case of 2,4,6-trichlorophenol either as a neutral molecule or as an ionized counterion species in association with the cetylpyridinium ion was also investigated in solution and at the silica/water interface.

Introduction Studies on the adsorption of surfactants at solid/water interfaces have been of continuous interest over many decades.1-5 It seems, however, that either some aspects have been neglected or their importance has been underestimated. In the case of ionic surfactants, one of these aspects is the consideration of the contribution of the counterions to the adsorption of the surfactant ion. Only hydrophilic surfaces will be discussed here. The influence of the type of counterion on the properties of micellar solutions, in particular the critical micelle concentration, the degree of counterion binding, or the aggregation number, is very well documented. The situation is far from being as well illustrated for adsorbed aggregates on surfaces. The classical work of Bijsterbosch6 on the condensation of the bromide ion on cationic aggregates on silica was published in 1974. It was shown then, among other important observations, that the adsorption of the counterion could only be observed when the surfactant experienced a certain degree of adsorption. This could be taken as evidence in favor of the formation, starting at a * To whom all correspondence should be addressed. E-mail: [email protected]. (1) Hough, D. B.; Rendall H. M. In Adsorption from solution at solid/ liquid interfaces; Parfitt, G. D., Rochester, C. H., Eds.; Academic Press: New York, 1983. (2) Ingram, B. T.; Ottewill R. H. In Cationic Surfactants; Rubingh, D. N., Holland, P. M., Eds.; Surfactant Science Series Vol. 37; Marcel Dekker: New York, 1991. (3) Furstenau, D. W.; Herrera-Urbina R. In Cationic Surfactants; Rubingh, D. N., Holland, P. M., Eds.; Surfactant Science Series Vol. 37; Marcel Dekker: New York, 1991. (4) Koopal, L. K. In Coagulation and flocculation; Dobias, B., Ed.; Surfactant Science Series Vol. 47; Marcel Dekker: New York, 1993. (5) Dobias, B. In Coagulation and flocculation; Dobias, B., Ed.; Surfactant Science Series Vol. 47; Marcel Dekker: New York, 1993.

given surfactant concentration, of aggregates whose surface charge density was large enough to induce the condensation of the counterion, much as in the case of ionic/nonionic surfactant mixtures where the counterion starts to condense on the aggregates above a critical composition of the mixed micelle.7,8 Other investigations have shown the large influence of the type of counterion on ionic surfactant adsorption at various solid/liquid interfaces.5 Rupprecht et al.9 observed in an extreme example of counterion effect taking tetradecylpyridinium as the surfactant ion, that the replacement of the chloride by the salicylate counterion increased the adsorption of the cationic surfactant at a silica/water interface by a factor of 6. The case of the salicylate counterion was very interesting but could be somewhat misleading because of the known specific interaction of this ion with pyridinium or ammonium headgroups, leading to the formation of viscoelastic solutions.10-13 Thus there was the need to study other salicylate derivatives which did not display this specific effect. Harwell and colleagues14,15 have carefully discussed a model to describe the adsorption of anionic surfactants on alumina where the statuses of the counterions of the inner (6) Bijsterbosch, B. H. J. Colloid Interface Sci. 1974, 47, 186. (7) Treiner, C.; Khodja, A. A.; Fromon, M. J. Colloid Interface Sci. 1989, 128, 416. (8) Treiner, C.; Mannebach, M. H. Colloid Polym. Sci. 1990, 268, 88. (9) Leimbach, J.; Sigg, J.; Rupprecht, H. Colloids Surf., A 1995, 94. (10) Wan, L. S. C. J. Pharm. Sci. 1966, 55, 1395. (11) Grasvholt, S. J. Colloid Interface Sci. 1976, 57, 575. (12) Underwood, A. L.; Anacker, E. W. J. Colloid Interface Sci. 1985, 106, 86. (13) Hoffmann, H.; Platz, G.; Rehage, H.; Schorr, W.; Ulbricht, W. Ber. Bunsen-Ges. Phys. Chem. 1981, 85, 255. (14) Hankins, N. P.; O’Haver, J. H.; Harwell, J. H. Ind. Eng. Chem. Res. 1996, 35, 2844. (15) Bitting, D.; Harwell, J. H. Langmuir 1987, 3, 500.

10.1021/la980890k CCC: $15.00 © 1998 American Chemical Society Published on Web 11/25/1998

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and the outer surfactant layers formed on hydrophilic surfaces were allowed to be different: the former being sensitive to the pH ionization of the solid surface and therefore to the pH of the solution, whereas the latter would be relatively insensitive to the local pH. It may be recalled that the inner surfactant layer contains more surfactant ions than there are ionic sites on the solid surface, so that small inorganic ions may be associated to the Stern layer for charge compensation. In that respect, some observations have been made recently16 which needed further investigation. It was shown that when a cationic surfactant such as cetylpyridinium chloride (CPC) is adsorbed at a silica/water interface, it may incorporate salicylate ions by an ionexchange mechanism. However, this salicylate uptake from solution decreases with increasing pH, although as it is well-known increasing the pH increases the adsorption of cationic surfactants. This conclusion was deduced both from distribution constant determinations16 and from calorimetric measurements.17 Assuming a surfactant bilayer at the adsorption isotherm plateau, these observations were interpreted as the consequence of the increasing ionization of the silica surface, which would repell a fraction of the salicylate of the inner surfactant layer into the solution. This interpretation implied that, at low pH, the counterions could interact with the surfactant headgroups of both surfactant layers but that as the surface ionization increased as the result of the rise in the pH, the counterions of the inner layer were repelled into the solution. It also implied that the pKa of the solutes played an important role in the coadsorption phenomenon: it was observed that when the CPC is adsorbed on a hydrophilic silica, the coadsorption of some weak acids increases with increasing pH below their pK values and decreases above the pK values. This could be due to an increasing coadsorption as the solute pK is approached from below and the ionic character of the molecule becomes increasingly important, an effect which would be counterbalanced by the increase of the ionization of the solid, as outlined above. This effect is reminiscent of that observed with ionizable molecules used in flotation experiments, such as dodecylamine, to which adsorption from solution at solid interfaces is dependent upon the pKa of the amine and the ionization of the solid surface.18 Finally, when the solute behaved predominantly as a neutral molecule, a partition model could be used to describe the interaction with the surfactant aggregates; however, as the solute became predominantly ionized above its pKa, a complexation model was necessary to represent the adsorption isotherm. The present investigation therefore has several goals. First, to remedy the deficiencies of the partition model for ionized solutes, a different experimental procedure and consequently a different analysis of the data from the one used previously16 for the determination of surfactant/ counterion interaction on a hydrophilic surface will be proposed. Second, to provide additional evidence for some of the above contentions, a wider variety of organic counterions will be investigated with additional techniques: 4-aminosalicylate, 5-aminosalicylate, and 2,4,6trichlorophenol. The two hydroxybenzoates bear a carboxy group in an ortho position, as in the case of the salicylate, ion but they do not induce a viscoelastic effect, most certainly because of the presence of the hydrophilic amino (16) Favoriti, P.; Mannebach, M. H.; Treiner, C. Langmuir 1996, 12, 4691. (17) Bury, R.; Treiner, C. Colloids Surf., A 1998, 139, 99. (18) Ananthpadmanabhan, K.; Somasundaran, P.; Healy, T. W. AIME Trans. 1978, 266, 2003.

Favoriti and Treiner Table 1. Characteristic Parameters for the Adsorption Isotherms of Three Cationic Surfactants at 1 % Silica/ Water (pH ) 4.4) and Air/Water Interfaces at 25 °C surfactant

cmca, mmol/L

Γmax (s/w), µmol/m2

As/w,a nm2/molecule

Aa/w,b nm2/molecule

CPC CP4-Sa CPSa

0.82 0.40 0.14

1.0 5.0 6.0

1.66 0.33 0.26

0.79 0.47 0.28

a In the absence of added salt. b Calculated from the Gibbs equation from surface tension experiments in the presence of 0.1 mol/L added salt with a common ion.

group. Third, we will extend this investigation to solid substrates other than silica such as titanium dioxide and alumina. Several approaches were employed: electrophoretic measurements, binding constant evaluations, surface tension data analysis, and considerations of the hydrophobicity of the solutes. Most of the data suggest that the main factor responsible for the large adsorption of cationic surfactants at solid/water interfaces when associated to the so-called hydrotropic counterion might be the possibility that these planar-shaped species which are highly concentrated at interfaces interact through a stacking mechanism. This stacking might decrease the overall Coulombic repulsion between surfactants of like charge, leaving the possibility for the cationic ions to come closer to each other. Materials and Methods Sodium salicylate (NaSa) was from Fluka (puriss pa), sodium 4-aminosalicylate (Na4-Sa) was from Aldrich (99% pure), 2,4,6trichlorophenol (TCP) was from Aldrich (98% pure), and 5-aminosalicylic acid (H5-Sa) was from Acros Organics (99% pure). These compounds were used as received. Cetylpyridinium chloride (CPC) was from Aldrich (98% pure). The preparation of cetylpyridinium salicylate (CPSa) and cetylpyridinium 4-aminosalicylate (CP4-Sa) was as follows. CPC (0.2 mol/L) was added to an equal concentration of either NaSa or Na4-Sa in pure ethanol. The solution was refluxed for 20 min. The solution was separated from the NaCl crystals and the supernatant evaporated to dryness under vacuum on a Rotavapor apparatus. The powder obtained was dissolved twice in chloroform to eliminate the remaining sodium chloride ions. Titration of the surfactant solution with silver nitrate showed the presence of chloride ion to be negligibly small. A further check of the purity of the surfactants was performed using the UV spectra of the surfactant and its organic counterion. The solubility of H5-Sa in water was determined at 25 °C and found equal to 6.0 × 10-3 mol/L. The pKa is equal to 3.2. At the pH of the adsorption measurements (see below), the acid was essentially used as the dissociated form. The critical micelle concentrations (cmc’s) of CPSa and of CP4Sa as well as their Krafft points were determined. The cmc’s at 25 °C were obtained from surface tension measurements (Kruss model K10T); they were equal respectively to 1.4 × 10-4 and 4.0 × 10-4 mol/L for CPSa and CP4-Sa. The surface areas in the presence of an excess of salt with a common ion were calculated from the Gibbs equation (Table 1). The Krafft points were determined visually as being equal respectivey to 13 and 19 °C. The cmc of CPSa was in agreement with literature values (cmc ) 1.5 × 10-4 mol/L).12 No value was found for CP4-Sa. The cmc of CPC was equal to 8.2 × 10-4 mol/L, again in agreement with literature values (Table 1). Most adsorption isotherms at solid/liquid interfaces were performed in the presence of 0.01 mol/L of sodium chloride except when stated otherwise. Silica, alumina, and titanium dioxide were very pure nonporous hydrophilic powders kindly provided by Degussa, France. Their BET surfaces as given by the manufacturer were the following: Aerosil 200 (200 ( 15 m2/g), Alumina C (100 ( 15 m2/g), and Titanium dioxide P25 (50 ( 15 m2/g). The P25 is a mixture of 70% anatase and 30% rutile. These compounds were made by

Binding of Anions to Cetylpyridinium Aggregates hydrolysis at high temperature of the corresponding chlorides. Purified water was used for all experiments. The experimental procedure for the adsorption isotherms was as follows: 0.2 g of solid was equilibrated with 20 mL of aqueous solution containing 1% solid substrate with sodium chloride and surfactant salt as needed at 25 °C. Some adsorption isotherms were also determined with 0.1% solid concentration in the presence of 5.0 × 10-4 mol/L NaCl. These experiments were needed for comparison purposes with the electrophoretic measurements which had to be performed at this low solid concentration. The equilibration times were 12 h for silica and 48 h for alumina and titanium dioxide. NaOH or HCl was used to adjust the solutions to the desired pH value. A combined glass electrode with a Tacussel (France) potentiometer was used for pH measurements. After ultracentrifugation (Sigma 2K15) at 25 °C, the supernatant was analyzed by UV spectroscopy (Cary 1E from Varian) at the same temperature. The maximum wavelengths used, the pH, and the molar extinction coefficients  (mol-1 cm-1) were as follows: NaSa (298 nm at pH ) 6;  ) 3600); Na4-Sa (299 nm at pH ) 7;  ) 7000); Na5-Sa (330 nm at pH ) 7;  ) 3500); TCP (289 nm at pH ) 9;  ) 4600); CPC (259 nm;  ) 3850). A number of experiments were also performed at constant solute concentration (e.g. NaSa), varying the surfactant concentration in order to investigate the uptake of the organic ion by the surfactant aggregates on the solid surfaces. Although the process was different from that described above, the experimental operations were conducted as described above. Finally ionexchange constants were also determined where the salt (e.g. NaSa) was added to the inhomogeneous solution at constant surfactant concentration. Here again, the experimental procedure was the same as that described above. Electrophoretic measurements were performed using a Zetasizer Model 4 (Malvern) with a fixed angle. Some additional measurements were also performed on a DELSA 440 (Coultronics), which allowed multiangular analysis. Degassed water was used for all experiments to avoid the formation of bubbles. The solid/volume ratio employed was 0.1% for experimental reasons with the present equipment. With the cells used in the presence of surfactant an added salt concentration (NaCl) of no less than 5 × 10-4 mol/L had to be used, as noted above. A more concentrated salt concentration would have displaced the organic counterions, Sa- and 4-Sa-, from the silica/water interface by the choride excess addition. The repeatability of the measurements was on the order of (0.5 mV. The mineral oxides’ isoelectric points (in the absence of surfactant) were determined in the presence of 0.01 mol/L NaCl. They were found at the following pH values: 2.9 (SiO2), 9.2 (Al2O3), and 6.3 (TiO2). The two instruments used gave, within experimental errors, the same numbers.

Results 1. Adsorption Isotherms at the Silica/Water Interface. Figure 1a presents the adsorption isotherms on a 0.1% silica dispersion of three surfactants: CPC, CPSa, and CP4-Sa at pH ) 5.4, without added salt. This is the spontaneous pH of that silica at 0.1% solid. Table 1 recalls some of the data pertaining to the three surfactants used in the present work. Note that the surface areas per molecule at the air/water interface, Aa/w, were obtained in the presence of 0.1 mol/L of added salt with a common ion because the Gibbs equation must be used in the presence of an excess of added salt. The As/w values were obtained from the classical equation with the implicit assumption of a surfactant monolayer at the solid/water interface. The results of Table 1 confirm the previous result obtained by Rupprecht et al.9 at a low pH value using tetradecylpyridinium salts adsorbed on a fused Silica, Aerosil OX50, similar to Aerosil 200, with a BET surface of 50 m2/g. The replacement of the chloride by the salicylate counterion considerably increases the adsorption of the cationic surfactant. In the latter case the surface areas for tetradecylpyridinium chloride and tetradecylpyridinium salicylate without added salt were equal respec-

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Figure 1. (a) Adsorption isotherm of cetylpyridinium salicylate (b), cetylpyridinium 4-aminosalicylate (x), and cetylpyridinium chloride (O) on 0.1% silica at pH ) 5.4 in the absence of added salt. (b) Variation of the electrophoretic mobility of 0.1% silica particles as a function of surfactant equibrium concentration (same symbols as in Figure 1a) in the presence of 5.0 × 10-4 mol/L added salt with a common ion.

tively to 1.84 and 0.33 nm2, which are to be compared to the present values of 1.66 and 0.26 nm2 for the corresponding cetylpyridinium derivatives. The difference between the adsorption plateaus for the structurally closely similar Sa- and 4-Sa- derivatives is small, being on the order of 20%. Thus, the Aa/w values at the air/water interface for the two salicylate derivatives are also much closer to each other than to that of the chloride ion. One should stress that the As/w value obtained at the silica/water interface in the absence of added salt is equal to the one obtained at the air/water interface in the presence of an excess of added salt. Insofar as it is assumed that surfactants may only form monolayers at an air/water interface, one is led to suggest that CPSa also forms a monolayer at the silica/water interface. In effect, wherever one finds at surfactant saturation a surface area per molecule at the air/water interface which is twice that of the solid/water interface, it is concluded that a bilayer is formed in the latter case. Numerous examples of that sort are given in the literature. For example, from neutron reflection experiments, the surface area of cetyltrimethylammonium bromide19 (CTAB), which forms a bilayer on an amorphous silica, is found to be equal to 0.30 ( 0.03 nm2. In the closely related case of decyltrimethylammonium bromide at the air/water interface,20 the same technique leads to a value of 0.58 ( 0.05 nm2. Both values have been obtained in the absence of added salt. This is an important point, because, classically, surface areas at air/water interfaces are deduced from the Gibbs equation. In the case of ionic surfactants, if this equation is to be applied in order to calculate the surfactant ion surface area, the addition of a salt with a common ion is necessary. Thus comparison (19) Lee, E. M.; Simister, E. A.; Thomas, R. K. Langmuir 1990, 6, 1031. (20) Lee, E. M.; Simister, E. A.; Penfold, J.; Ward, R. C. J. Phys. Chem. 1989, 93, 381.

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Table 2. Surface Areas of Cetylpyridinium Chloride in the Presence of Added NaCl at 1% Silica/Water and Air/Water Interfaces at 25 °C C(NaCl), mol/L

Γmax (s/w), µmol/m2

As/w, nm2

Aa/w,a nm2

0.01 0.05 0.10 0.20 0.01 0.01

2.2 3.4 3.8 4.5 3.9 4.6

0.75 0.49 0.44 0.37 0.42 0.36

0.79 0.82 0.78 0.81

pH 4.2

6.5 8.5

a Calculated from the Gibbs equation using previously published surface tension data: Monticone, V., Treiner, C. Colloids Surf., A 1995, 104, 285.

is not always straightforward between solid/water and air/water surface area data. Table 2 shows some results obtained for CPC at both interfaces. It can be seen that addition of salt does not change the surface area at the air/water interface above a NaCl concentration of 0.01 mol/L, whereas, at the silica/ water interface, addition of salt leads to the formation of a bilayer and hence a factor of about 2 between both sets of measurements at higher salt concentration. This is not a surprising result: the Aa/w values of sodium alkylbenzenesulfonate derivatives were likewise found to be constant from 9 × 10-4 to 7 × 10-2 mol/L.21 It is interesting to compare the partition coefficient Poct of salicylic acid and that of 4-aminosalicylic acid in the two-phase octanol/water binary.22 This system has often been taken as an index of hydrophobicity for undissociated solutes. The log Poct values of the two salicylates are respectively 2.26 and 0.89. Thus, as measured on this scale, the hydrophobicities of the two acids are substantially different. Therefore, this parameter alone would not explain the close values of the adsorption plateaus at the silica/water interface, as displayed by the surface area per molecule for the salicylate derivatives of Table 1. Note that log Poct for the undissociated form of 2,4,6-TCP is equal to 3.69. 2. Electrophoretic Measurements. Figure 1a presents a comparison, as classically displayed,23-25 between the adsorption isotherms and the electrophoretic mobility of the particles. Recall that both sets of measurements were performed for 0.1% silica dispersion at pH ) 5.4 and 5.0 × 10-4 mol/L NaCl. As is often observed, there is a close parallelism between the profiles of the two series of curves with two interesting features: (i) Despite the very large differences between the plateau values of the adsorption isotherms of CPC, CPSa, and CP4-Sa on the silica/water interface, there is no difference, within experimental error, between the maximum electrophoretic velocities of the dispersions. (ii) The point of zero charge is attained with much less adsorbed surfactant for CPC than for CP4-Sa or CPSa. Another way of presenting this result is displayed in Figure 2, where the variation of the ζ potential as a function of the concentration of adsorbed surfactant is displayed. The quantities of adsorbed surfactant necessary to obtain the point of zero charge were deduced from these curves. One obtains the following values for, respectively, CPC, CP4(21) Lascaux, M. P.; Dusart, O.; Granet, R.; Piekasrki, S. J. Chim. Phys. 1983, 80, 615. (22) Hansch, C.; Leo, A. Substituent constants for correlation analysis in chemistry and biology; Pomona College: Claremont, California, 1979. (23) Somasundaran, P.; Fuerstenau, D. W. J. Phys. Chem. 1966, 70, 90. (24) Thomas, F.; Bottero, J. Y.; Cases, J. M. Colloids Surf., A 1989, 37, 269. (25) Esumi, K.; Shibayama, M.; Meguro, K. Langmuir 1990, 6, 826.

Figure 2. Variation of the ζ potential of 0.1% silica particles as a function of adsorbed surfactant concentration (same symbols as in Figure 1).

Figure 3. Variation of the salicylate ion coadsorption on 1% silica as a function of cetylpyridinium chloride equilibrium concentration in the presence of 0.01 mol/L NaCl: constant total sodium salicylate concentration, 4.0 × 10-4 mol/L; pH ) 4.4 (b); pH ) 6.5 (O); pH ) 8.5 (2).

Sa, and CPSa at pH ) 5.4: 0.10, 0.45, and 0.85 µmol/m2 with an estimated error on the order of (0.05 µmol/m2. Again, the electrophoretic results point out the similarities of the behavior of the surfactants with organic counterions when compared to that of those with a chloride counterion. 3. Complexation Constants for Counterion Effects. In the case of ionic surfactants, the strength of the interaction between the counterion and the surfactant ion is indirectly indicated by the value of the degree of condensation of the counterions at the cmc. However this interaction may also be evaluated as suggested recently16 by considering the counterion as an additive. Keeping the counterion constant and varying the surfactant concentration, one can observe first the uptake of the anions of interest by the surfactant aggregates until the cmc is attained. Then, above the cmc, as micelles are formed, the anion under study will be distributed between the surface aggregates and the free micelles in solution. The results of this procedure are shown in Figure 3 for SaNa at a constant concentration of 4.0 × 10-4 mol/L with varying CPC concentrations on Aerosil 200 at three different pH values in the presence of 0.01 mol/L NaCl. At a higher additive concentration, the salicylate ions will be completely desorbed from the solid/liquid interface. Experiments were also performed at several constant NaSa concentrations from 2.0 × 10-4 to 8.0 × 10-4 mol/L. It was observed that the adsorption of CPC was slightly increased by the presence of the Sa- ions at low surfactant

Binding of Anions to Cetylpyridinium Aggregates

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Table 3. Calculated Values of Complexation Constants of Ionic Solutes Coadsorbed with Cetylpyridinium Aggregates at Solid/Water Interfaces or on Free Micelles in the Presence of NaCl at 0.01 mol/L (Molality Scale) systems

pH

Γ(cpc), ×104 mol/g

Γmax(sol),c ×104 mol/g

Kb, L/mol

KM, L/mol

SiO2/SaSiO2/SaSiO2/SaSiO2/4-SaSiO2/5-SaSiO2/TCP

4.4 6.5 8.5 6.5 6.5 8.5 3.6 9.1 10.2

4.4 7.5 9.0 7.5 7.5 8.5 4.6 2.0 2.4

3.30 5.25 5.13 3.21 4.62 5.31 3.6 1.44 1.68

27000 20000 12800 1600 2100 86000 6200b 22000 7000a

8500 8800 8000 700

TiO2/SaAl2O3/Sa-

30000 3300b 12000 9000

a Approximated value due to the salicylate ion adsorption on alumina in the absence of surfactant (see text). b Calculated partition coefficient values of the adsorption constant Pads and of the micellar solubilization constant Pmic for the neutral TCP molecule (see text). c Γmax(sol) is the maximum quantity of added solute used for the determination of the isotherm at a constant cpc concentration Γ(cpc) adsorbed at the solid/water interfaces.

coverage but was not modified above a surfactant coverage on the order of 20%. These experiments are the analogue of the adsolubilization experiment, sometimes also called coadsorption, used by several authors for the uptake of neutral molecules from a solution.27-30 The interpretation of the profiles of the coadsorption curves is as follows. The salicylate ion does not adsorb on silica in the absence of CPC. As CPC adsorbs, forming aggregates (admicelles) on the solid particles, Sa- ions coadsorb on these aggregates by an ion-exchange mechanism with the chloride ion, up to the cmc. As the CPC concentration exceeds the cmc and micelles are formed in solution, the Sa- ions are distributed between the adsorbed aggregates and the free micelles. Ultimately, at a CPC concentration above 0.1 mol/L, 95% of the Sa- ions are condensed onto the free micelles, leaving few ions on the modified silica surface. Below the cmc, the coadsorption of the salicylate ions depends on the pH of the systems, as outlined previously. Note that the decrease is the same at the three pH values although the maximum amount of surfactant adsorbed at the plateau differs significantly with the change in pH of the system (Table 3). The interaction between the salicylate ion and CPC on silica had been determined previously both with adsorbed aggregates and with free micelles. Three main conclusions had been drawn:16 (1) The CP+/Sa- interaction is larger for the former aggregates than for the latter. (2) In the case of adsorbed aggregates, this interaction decreased with increasing pH. (3) At higher pH, the CP+/Sainteraction approaches that observed with free micelles. The formalism which was adopted was that of a distribution model16 which made use of a solute partition coefficient between aggregates and water whatever the type of aggregates: adsorbed aggregates on the particle surface or free micelles in solution. This approach was satisfactory in the case of neutral molecules. However when strong interactions occurred between solute and (26) Schieder, D.; Dobias, B.; Kumpp, E.; Schwuger, M. J. Colloids Surf., A 1994, 88, 103. (27) Monticone, V.; Mannebach, M. H.; Treiner, C. Langmuir 1994, 10, 2395. (28) Bernard, R.; Fuchs, E.; Strnadova, M.; Sigg, J.; Vitzhum, J.; Rupprecht, H. Prog. Colloid Polym. Sci. 1990, 83, 110. (29) Lee, C.; Yeskie, M. A.; Harwell, J. H.; O’Rear, E. A. Langmuir 1990, 6, 1758. (30) Harwell, J. H.; Hoskins, J. C.; Schechter, R. S.; Wade, W. H. Langmuir 1985, 1, 251.

Figure 4. Coadsorption isotherm of the salicylate ion on 1% silica: pH ) 4.4 (b); pH ) 6.5 (O); pH ) 8.5 (2); NaCl concentration, 0.01 mol/L; constant total cetylpyridinium chloride concentration, 7.0 × 10-3 mol/L.

Figure 5. Coadsorption isotherms of sodium salicylate (b), 4-aminosalicylate (O), and 5-aminosalicylate (2): 1% silica; pH ) 6.5; constant total cetylpyridinium chloride concentration, 7.0 × 10-3 mol/L; NaCl concentration, 0.01 mol/L.

surfactant ions, the distribution model did not fit the experimental data at very low surfactant concentration. Therefore a different experimental procedure was adopted here and a more rigorous model was applied and extended to the other ions such as TCP and Na4-Sa as well as to NaSa coadsorbed with CPC on TiO2 and on Al2O3. 3.1. Coadsorption of Added Salt on Adsorbed and Free Aggregates. The present adsorption experiments were performed in all cases at a constant CPC concentration which was chosen to be just below the equilibrium cmc, varying the solute (e.g., NaSa) concentration. This procedure will be shown to be more rigorous than the one described previously at constant solute (e.g. NaSa) concentration. It had the advantage of representing more clearly the coadsorption effect below and above the cmc. Figure 4 presents the profiles obtained with NaSa at the three pH values 4.4, 6.5, and 8.5. Note that the CPC concentrations differ at each pH value because the CPC adsorption on silica increases with the pH in the case of silica (see Table 3). Figure 5 presents the results obtained at pH ) 6.5 for the three sodium salicylate salts NaSa, Na4-Sa, and Na5-Sa. Here the same initial surfactant concentration was used for the three salts. Such profiles could be represented formally by a Langmuir-type isotherm, assuming a 1 to 1 complexation model in a solute surface coverage up to 0.5 or 0.7, depending upon the system analyzed. The Langmuir equation may be written as

θad ) KbCeq 1 - θad

(1)

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where θad ) Cad/Cad,max is the ratio of the concentration of coadsorbed salicylate ions for a given CPC concentration to the maximum salicylate coadsorption. Kb is the apparent association constant, and Ceq is the concentration of free added ions. Rearrangement of eq 1 leads to the well-known form

1 1 1 ) + Cad Cad,max KbCad,maxCeq

(2)

Thus, if the Langmuir formalism may be applied here, the change of 1/Cad with 1/Ceq should give a straight line from which the corresponding association constant is evaluated. The decreasing part of the curve of Figure 3 may be represented by the combined Langmuir equations (eqs 1 and 3). Equation 3 is the equivalent of eq 1 for micellar aggregates above the cmc.

θmic ) KmicCeq 1 - θmic

(3)

with the condition that the total organic ion concentration is

Ct ) Cad + Ceq + Cmic

(4)

The same definition is applied to θmic and to θad. A number of reasonable assumptions have to be made: (a) Cmic,max ) pDmic, where Dmic is the concentration of micellized surfactant and p is the number of sites per micellized surfactant (p is taken as proportional to the degree of counterion condensation on micelles (see Table 1)). (b) The added ion concentrations are small enough so that the following conditions hold: θad , 1 and θmic , 1. One gets finally in terms of the concentration of coadsorbed counterions on micelles

Cad )

Ct Kmic p(Ccpc,eq - cmc) 1+ KbCad,max

(5)

p has been shown to vary from 0.66 to 0.88 using results on polychlorophenols in CTAB31 micellar solutions. The value 0.8 has been adopted here for Sa- on micelles, and 0.5 has been adopted for 4-Sa-. With Ct and Cad,max being known from experiment and Kb from the coadsorption analysis outlined above, Kmic is the only unknown parameter in eq 5. This equation is the equivalent of that derived previously for nonionic solutes using a partition coefficient formalism. Figure 6 shows an illustration of eq 2 for the determination of Kb for the three different salicylate derivatives. Table 3 presents the results obtained for all solutes at different pH values and on the various substrates. A number of observations can be made. First, the constancy of Kmic at each pH value from 4.4 to 8.5 is remarkable. It corresponds of course to the profiles observed in Figure 3 where the desorption curves coincide. This shows that the analysis has been correctly performed, as the coadsorption of salicylate ions on free micelles should be independent of the pH of the system. Second, one observes that Kb is systematically larger than KM by a factor of 2-3. The only exception is for the salicylate ion adsorbed on alumina at a low pH value. The salicylate ion (31) Berlotti, S. G.; Garcia, N. A.; Gsponer, H. E. J. Colloid Interface Sci. 1989, 129, 406.

Figure 6. Representation of eq 5 for the salicylate ion (b), 4-aminosalicylate (O), and 5-aminosalicylate (2) (same conditions as those in Figure 4).

slightly adsorbs onto alumina32 (see below). All other solutes studied did not adsorb on the substrates in the absence of surfactant. The adsorption of the organic ions directly on the alumina surface evidently complicates the analysis. We therefore confirm earlier results16 obtained showing that, contrary to simple electrostatic considerations, the salicylate ion adsorbs on alumina although, at the pH of the experiments, alumina and salicylate bear the same electric negative charge. This point will be considered further in the Discussion. 3.2. Ion-Exchange Constants. It was interesting to consider the effect of the presence of NaCl on the coadsorption results. One of the reason for the presence of added salt with a common ion was to enhance the adsorption of the surfactant. As a consequence the association of the added organic ions with the systems implied the replacement of the chloride ions by the organic ones. This effect could be evaluated by ion-exchange experiments. A simple procedure was adopted where NaCl was added to the CPSa/silica or CP4-Sa/silica systems. The release of the organic ions upon this addition was monitored. The classical expression used for ion-exchange resins, for micellar catalysis, or for the flotation of minerals may be written

Kab )

(A)eq(B)ad (A)ad(B)eq

(6)

Figure 7 displays the variation of (Cad/Ceq)X- as a function of (Cad/Ceq)Cl-. The following constants were obtained from the slopes of these lines:

Kab ) 250 ( 50 for the exchange Sa-/ClKab ) 13 ( 3 for the exchange 4-Sa-/ClSuch values may be compared to that obtained by Thalody and Warr33 at the air/water interface saturated with tetradecylammonium bromide: Kab was found equal to 49 ( 10 for the Sa-/Br- exchange. One can deduce from such results that the condensation of the organic ions on the cationic aggregates is strongly favored at solid/water or at air/water interfaces. (32) Thomas, F.; Bottero, J. Y.; Cases, J. M. Colloids Surf., A 1989, 37, 281. (33) Thalody, B. P.; Warr, G. G. J. Colloid Interface Sci. 1995, 175, 297.

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Figure 8. Surface tension of cetylpyridinium chloride as a function of concentration in the presence of a constant concentration of 4.0 × 10-4 mol/L added salts. Figure 7. Determination of ion-exchange constants for the binary systems (eq 6): salicylate/chloride (b) and 4-aminosalicylate/chloride (O).

Discussion The Discussion will first address the electrophoretic mobility and the adsorption iostherms obtained for the cationic surfactants with organic counterions; in the second part, the effect of the organic ions considered as additives will be discussed. 1. Surfactants with Organic Counterions. The close relationship displayed by the profiles of surfactant adsorption and electrophoretic mobility curves has been noted by a number of authors. However, some observations by Bitting and Harwell14 have raised questions about the interpretation of such relations. In the present investigation, the following observations can be made: (i) The adsorption of surfactants may increase with increasing concentration even as the electrophoretic mobility remains constant. The same observation had been made for the adsorption of dodecyl sulfate with various alkali counterions on alumina. Furthermore it is interesting to point out that the surfactant concentrations for which the ζ potential is nil (respectively, 0.10, 0.45, and 0.85 µmol/m2 for CPC, CP4-Sa, and CPSa) are in (about) the same relative ratios as the maximum adsorption plateaus (1.0, 5.0, and 6.0 µmol/m2). (ii) The particle electrophoretic mobility depends strongly upon the type of counterion in the low-surfactantconcentration regime, but not at the plateau values. Interpretation of these results at high charge density may be misleading, as noted by Hough and Rendall1 in their classical review paper. However, electrophoretic measurements provide reliable information on the variation of the Stern layer at low concentrations, that is, near the isoelectric point. Esumi et al.35 have observed in the case of a series of cationic surfactants bearing one to three hydrocarbon chains of identical lengths adsorbed on a hydrophilic silica that the electrophoretic mobilities at the point of zero charge (pzc) are independent of the number of alkyl chains on the surfactant ion. In the general case cationic surfactants display a high dependence of the pzc upon the alkyl chain length and the type of headgroup. Unfortunately in both of these studies the particle velocities at surfactant saturation were not attained, which prevents a complete comparison with the present observation. The variation of particle mobility observed here resembles somewhat that described by Dobias at lower (34) Dao, K.; Bee, A.; Treiner, C. J. Colloid Interface Sci. 1998, 204, 61. (35) Esumi, K.; Matoba, M.; Tamanaka, Y. Langmuir 1996, 12, 2130.

surfactant concentrations in the case of CTAB with added ethylenediamine tetraacetic acid (EDTA):5 the particle mobility was much lower in the presence than in the absence of EDTA. This behavior was interpreted as the consequence of complex formation between EDTA and CTAB. Although the complexation phenomenon is real, it may not be the only source of the observed behavior. There is hardly any complexation between the aminosalicylate ion and the pyridinium moiety; however, the electrophoretic behavior of this counterion is very close to that of the salicylate ion. (iii) The Kb values of Table 3 confirm our earlier suggestion of a strong binding of the Sa- ion with the pyridinium moiety. It also shows that the binding is much weaker with the two other salicylates. The difference in the ion-exchange constants also indicates a much stronger interaction of the Sa- than 4-Sa- ions with the pyridinium headgroup. Note, however, that the maximum amount of organic counterion which may be incorporated in the adsorbed surfactant layers is about the same for 5-Saand Sa-. The type of surfactant/ion interaction involved here may be characterized with the help of surface tension measurements. If the surface activity of the alleged complex formed is higher than that of the surfactant molecule, a surface tension minimum may be observed close to the cmc. This behavior was observed for example when a surfactant dimer was formed below the cmc18 or when a specific interaction occurred, such as in the case of ferric ions in the presence of sodium octylbenzenesulfonate.34 Figure 8 presents the surface tension of CPC as a function of concentration in the presence of 4.0 × 10-4 mol/L NaSa, Na4-Sa, and TCP (at pH ) 9.1, where the phenol is completely ionized). No minimum is observed in any of these cases. This shows that though the interaction between CP+ and the various anions is indeed strong (as discussed below), it does not induce the formation of a surface active complex. These various observations may be summarized in the following manner: (i) The surfactant isotherms at the silica/water and the air/water interfaces show that 4-Saand Sa- display similar behaviors. (ii) The electrophoretic results also show that 4-Sa- and Sa- behave similarly. (iii) The binding isotherms show a strong preferential interaction of the pyridinium moiety with the Sa- ion but not with the 4-Sa- or the 5-Sa- ions. (iv) The hydrophobicity of the corresponding acids suggests again a large difference between the Sa- and 4-Sa- ions. The conclusion one may infer from such considerations is that the strong binding displayed by the salicylate ion with respect to the surfactant headgroup does not explain

7500 Langmuir, Vol. 14, No. 26, 1998

the amplitude of the adsorption isotherms. Nor does the hydrophobicity of the organic counterions. One is therefore led to suggest that the observed adsorption behavior is related to some other property of the investigated counterions. One ad-hoc hypothesis will be put forward below which is related to a property shared by many salicylate derivatives in aqueous solutions: it is the stacking effect. It is known that this effect concerns planar hydrophobic molecules such as bile salts or aromatic molecules which associate to form small aggregates in water by a mechanism which is distinct from the classical hydrophobic effect and is usually observed at relatively high solute concentration. The surface areas presented in Table 2 show that the values of As/w for CPC decrease asymptotically from 0.75 nm2 with 0.01 mol/L NaCl down to 0.37 nm2 at 0.2 mol/L added NaCl, whereas, at the air/water interface, the value of Aa/w remains constant and equal to 0.80 nm2 in the same salt concentration range. The latter value is therefore the minimum one which can be obtained for a monolayer at a maximum screening effect at the air/water interface. As a consequence, the minimum value obtained at the silica/water interface (which is about half that deduced from the surface tension data) must correspond to that of a nearly complete bilayer, with a surface area per surfactant molecule which remains constant and equal to the value found at the air/water interface. In the case of CP4-Sa and CPSa, the As/w values are equal respectively to 0.33 and 0.26 nm2 at the silica/water interface in the absence of added salt, that is, without the need for a screening effect. These values are very close to the minimum values of the cross section of a surfactant hydrophobic tail. In the case of 1-octanol, with no repulsion between headgroups, one gets a Aa/w value of about 0.24 nm2. Incidently, this means that the value of the adsorption plateau obtained with the salicylate counterion and the cetylpyridinium surfactant ion should be the maximum surfactant adsorption which can be obtained for this hydrocarbon chain length at this silica/water interface. One is led to the following conclusions: (i) Comparison of the surfactant surface areas at the air/water and at the silica/water interfaces strongly suggests that CPSa may form a monolayer at the silica/ water interface at complete surface coverage. The pyridinium headgroups should be pointing toward the solid surface. The reason for the latter suggestion is that, by increasing the pH, and hence the surface charge density, the Kb values of the salicylate/pyridinium system decrease because then a fraction of the complexed salicylate ions are released into the solution. This effect is the consequence of an increasing repulsion between the surface ionic sites and the organic counterions, which in turns implies some contact between the ionic sites on the solid and the salicylate ions. The repulsion effect would not take place if the organic counterion pointed with the surfactant headgroup toward the bulk of the solution. This point can be looked upon as a consequence of the model proposed by Hankins et al.,14 where only the conterions of the inner layer are pH dependent. Evidently, the picture of a CPSa monolayer at the silica/ water interface is somewhat disturbing because it implies that the surfactant hydrocarbon tails are pointing toward the aqueous solution. This apparent contradiction cannot be resolved by thermodynamics alone. (ii) Essentially the same conclusion should be derived from the present data for the aminosalicylates which show a behavior similar to that of the salicylate counterion as far as adsorption at interfaces and electrophoretic profiles are concerned. The values of Aa/w and As/w for CP4-Sa are

Favoriti and Treiner

too close to each other to suggest the need for a complete bilayer. A model of a partial bilayer would be more appropriate. These apparently conflicting observations may be reconciled if one recalls one of the specific properties of salicylate derivatives: the possibility of approaching each other though a stacking effect. It has been known for a long time that hydrotropic salts may solubilize scarcely soluble molecules at high salt concentrations as the result of the stacking of mostly planar ions at high concentration. Sodium salicylate belongs to this class of so-called protosurfactants.36 The following hypothesis may be put forward: counterions at interfaces are very concentrated species. This is true for the Sa- and the 4-Sa- ions at both the silica/water and the air/water interfaces. There may therefore be the possibility that these ions may interact through this stacking mechanism.36-40 Such interactions and the solute solubilization which follows need rather large salt concentrations in bulk aqueous solutions. However, these concentrations are attained readily at interfaces. If the counterions, although of like charge, may interact favorably with each other, the surfactant hydrocarbon tails might come closer to each other as a fraction of the surfactant headgroups is anyway partly screened by the surface ionic sites. This mechanism would enable the solid surface to accommodate more surfactant ions. The somewhat lower adsorption obtained with CP4-Sa as compared to CPSa would be due to the less effective stacking obtained with an amine moiety in the para position. It must be emphasized that the solubilization mechanism which is at work with protosurfactants such as sodium salicylate is often linked to a complexation effect because such systems are used in pharmaceutical formulations where the drugs are themselves ionic species.39 This is by no means a general explanation, as the large increase of solubilization of n-octane in water by the addition of sodium salicylate derivatives illustrates conclusively.36 Sodium salicylate has also been used with oil and water for the formation of microemulsions.40 Again, whether a surfactant bilayer can be formed in the case of these organic counterions at the silica/water interface should be the subject of closer scrutiny using appropriate physicochemical techniques.19,20 (iii) We tend to conclude that the larger Aa/w values obtained by CPC at the air/water interface at all added NaCl concentrations compared to that at the silica/water interface at low salt concentration may be due in part to the presence of the chloride counterions between the headgroups of the inner surfactant layer. As recalled above, the apparent decrease of the CPC surface area with added salt in the case of a CPC at the hydrophilic/water interface is due to the formation of the bilayer. Note that the value of Aa/w for the CP4-Sa derivative is not much higher than that for CPSa although the Kb value of the former counterion is hardly a tenth of the value of the latter. Thus, for the sake of self-consistency, there is no need to admit the formation of a bilayer in that case either. One is led to suggest that the complexation phenomenon plays a minor role in the adsorption behavior of the (36) Ho, P. C. J. Chem. Eng. Data 1985, 30, 88. (37) Darwish, I. A.; Florence, A. T.; Saley, A. M. J. Pharm. Sci. 1989, 78, 577. (38) Badwan, A. A.; El-Khordagui, L. K.; Saley, A. M.; Khalil, S. A. Int. J. Pharm. 1983, 13, 67. (39) Attwood, D.; Florence, A. T. Surfactant Systems; Chapman and Hall: London, 1983; p 370. (40) . Ho, P. C. J. Phys. Chem. 1981, 85, 1445.

Binding of Anions to Cetylpyridinium Aggregates

Langmuir, Vol. 14, No. 26, 1998 7501

Table 4. Standard Thermodynamic Functions for the Binding of Sodium Salicylate and Sodium 4-Aminosalicylate to CPC Aggregates Adsorbed at Various Solid/Water Interfaces at 25 °C (Molar Scale, See Text) systems

pH

∆Gb°, kJ/mol

∆Hb°,a kJ/mol

∆Sb°, J/(deg/mol)

SiO2/SaSiO2/SaSiO2/4-SaTiO2/Sa-

4.4 8.5 6.5 9.1

-30.0 -27.8 -23.0 -29.5

-24.1 -18.5 -15.8 -17.2

19.8 31.2 24.1 24.1

a

Data from ref 17.

surfactant and that it is the stacking effect induced by the organic counterions studied which is the most important parameter. It is suggested also that the presence of the small inorganic ions near the solid surface sterically hinders the cationic surfactant ions, preventing them from coming into close contact. In their adsorption model, Hankins et al.14 consider that the smaller alkali counterions (in the case of anionic surfactants adsorbed on alumina) are retained in the inner surfactant layer by complexation with the positive sites at the alumina/water interface, an effect which shields the headgroup repulsions. The same model could, in principle, be applied to the present system although a stabilization by complexation with halide ions on silica seems unlikely. Goloub et al.41 also pointed out that, at high ionic strength, it becomes more difficult for the cationic surfactant headgroup to approach the solid surface as more inorganic counterions are crowded at the silica/water interface. Finally, previous data of the standard enthalpy of binding for some of the systems of Table 3 were used in conjunction with the present binding constants to evaluate the standard entropy changes associated with the present process of exchange between chloride and salicylate ions at the solid/water interface. Table 4 presents the results of the calculations. As the Kb values of Table 3 are expressed in the molal scale, the association constants were translated into a molar scale using a partial molar volume of 0.15 L/mol, as done previously.15 Although this procedure is somewhat arbitrary, it enables one to obtain at least a semiquantitative evaluation of the standard entropy changes of Table 4. They are all positive. Positive entropy changes have also been obtained using the micellar KM values and the corresponding standard enthalpies of interaction at the micellar surface (not shown). For the same process at the air/water interface, Thalody and Warr33 find a negative change, thus ascribing the chloride/ salicylate exchange to an enthalpy-driven process. However, interpretation of entropy changes in such complicated chemical systems without a theoretical model as a guideline may be misleading. It will not be attempted at this stage. 2. Other Considerations. The results of Table 3 show that when the solid surface is negatively charged, complex formation occurs between the salicyate ion and the pyridinium group, whatever the nature of the mineral oxide. Hence the value of Kb for Sa- on TiO2 at pH ) 9.1 (Kb ) 22 000) is equal within experimental error to the value obtained on silica at pH ) 6.5 (Kb) 27 000). Considering the isoelectric point of these two mineral oxides, one may note that the difference (pH - iep) is almost the same for the two solids, being respectively equal to 2.8 and 3.6. One may conclude that the nature of the (41) Goloub, T. P.; Koopal, K. L.; Bijsterbosch, B. H. Langmuir 1996, 12, 3188.

Figure 9. Variation of TCP coadsorption (total concentration, 4.0 × 10-4 mol/L) as a function of cetylpyridinium chloride equilibrium concentration: pH ) 8.5 (O); pH ) 3.6 (b).

solid surface in surfactant/counterion interactions intervenes only through the amount of surfactant the solid allows to adsorb on its surface. As noted previously, the case of CPSa on Al2O3 is different. The Kb value is substantially lower than that in the other cases. One reason could be that the pH at which the experiments were performed (10.2) is much closer to the solid iep (9.2) than it is with the two other oxides. Experiments with such systems at pH values above the limit 10-11 may introduce systematic errors. However, another complication arises with this system: there is a clear adsorption of the salicylate ion on alumina above the iep of that oxide, that is, in a pH domain where the solid is negatively charged in the absence of added surfactant. This observation has been made and analyzed recently.32 Thus, the value of Kb for this system is too small because the formal method used to extract the equilibrium constants did not allow for some initial ion adsorption on the solid surface in the absence of surfactant. Note, however, that the value of KM is equal within experimental error to that obtained in the presence of the silica dispersion. This is only natural, as the KM values should indeed by independent of the nature of the solid phase. Finally, the case of TCP is of special interest. This phenol was studied at low and at high pH values at the silica/ water interface. As the pK of TCP is equal to 6.0, this solute is undissociated at pH ) 3.6 and completely ionized at pH ) 8.5. Figure 9 presents first the coadsorption behavior of TCP at a constant solute concentration of 4.0 × 10-4 mol/L at the two pH values as a function of the CPC equilibrium concentration. This system displays essentially the same characteristics that were shown by previously studied systems with some differences. As a neutral solute, the profile is classical in this representation: coadsorption below the cmc as the surface coverage by CPC is increased and desorption above the cmc as free micelles are formed in the solution. The case of the ionized solute is somewhat different. Although the ionized TCP does not adsorb on the silica surface in the absence of surfactant, a very small amount of adsorbed CPC is enough for the total incorporation of the TCP ions (at the concentration used of 4.0 × 10-4 mol/L) to the silica/ water interface. The TCP adsorption remains constant until free micelles are formed above the cmc: the coadsorption then decreases, as in all other cases studied. Note that, at both pH values, the total initial concentration of the phenol is incorporated to the interface; only the concentration of CPC needed for complete ion uptake is

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based upon the strong binding of this ion to the pyridinium group. It should not be easily applied to organic counterions which do not present the strong complexation effect outlined above. Another interpretation of the difference in the strength of the binding of ions to free micelles and adsorbed aggregates would take advantage of the fact that micelles possess a high curvature radius when compared to that of adsorbed bilayer aggregates. Calculations based upon electrostatic considerations42 alone have shown that the degree of counterion binding should be higher on flat surfaces than on curved surfaces. The observations described here and previously may be the consequence of such an effect. Figure 10. Representation of eq 5 for the salicylate ion at pH ) 8.5 (O) and for the TCP ion at pH ) 9.1 (b).

different for the ionized and the unionized species: minute amounts at a high pH value and an equilibrium surfactant concentration close to the cmc at low pH. It is for this solute that the largest ion complexation constants have been obtained in the present investigation either on the adsorbed CPC aggregates or on the free micelles in solution. This result is related to the higher hydrophobicity of TCP as evaluated by the log Poct value of the neutral species. The data were analyzed at pH ) 8.5 using eq 5 for the completely dissociated form (Figure 10). At pH ) 3.8 the calculation of Pads corresponding to the coadsorption process (below the cmc) was performed using the data at constant solute concentration and the previously described distribution model for undissociated chemical species. Above the cmc, the desorption coefficient PM, which is in fact a solubilization constant, was calculated as described previously for undissociated species from the decreasing portion of the curve of Figure 8. The partition coefficient Pads and the PM obtained cannot be directly compared to the Kb values obtained for the dissociated form of TCP. The corresponding results are therefore noted specifically in Table 3. Note finally that the ratio of about 3 which has been observed for weak acids between the ion complexation constants on the adsorbed aggregates and on free micelles16 is again found here as with the salicylate derivatives and for TCP. The origin of this extra coadsorption effect has been discussed before. Hankins et al.14 have described a somewhat similar effect in the case of alkali counterions with dodecyl sulfate on alumina. Their calculations indicated that the origin of this difference may be the different status of the counterions, whether they belong to the inner or to the outer layer of the surfactant aggregates (the admicelles). In the case of the salicylate ion, it was recently suggested16 that the two-status model could explain the decrease of the thermodynamics of ion complexation with increasing pH on silica. However, the arguments were

Conclusions Analysis of the binding constants of organic ions to cationic surfactants at the solid/water interface, electrophoretic measurements, analysis of surface area data, and adsorption isotherms of the surfactants with organic counterions have been performed. It was observed that some organic counterions belonging to the so-called hydrotropic series, when associated with cationic surfactants, favor the adsorption of the surfactant ions on hydrophilic solids in water. This effect may be the consequence of a stacking of a fraction of these planar aromatic counterions at air/water or solid/water interfaces, which brings closer together the cationic surfactant ions. The specific interaction which takes place between the salicylate ion and the pyridinium group favors the uptake of the salicylate ion from solution at low surfactant concentration but plays a minor role at high surfactant concentration, where the interaction between the counterions through the stacking effect favors surfactant adsorption. An analysis of the surface area per molecule data at the silica/water and the air/water interfaces suggests that, with these hydrotropic counterions, cationic surfactants might form only monolayer structures at the silica/water interface. The formation of surfactant bilayers would be associated with the presence of small hydrophilic counterions, a fraction of which are located near the solid surface, preventing adjacent surfactant headgroups from approaching each other. The binding constants of the salicylate ion with CPC on hydrophilic titanium dioxide and alumina do not contradict this analysis, although, in the case of alumina, the adsorption of the salicylate ion on the negatively charged alumina may complicate the formal analysis of the data. Acknowledgment. We are grateful to the Laboratoire de Chimie Macromole´culaire of the University Pierre et Marie Curie for making the Zetasizer available to us. LA980890K (42) Gunnarsson, G.; Jonsson, B.; Wennerstrom, H. J. Phys. Chem. 1980, 84, 3114.