11867
2007, 111, 11867-11869 Published on Web 09/26/2007
The Chameleon-like Nature of Zwitterionic Micelles. Control of Anion and Cation Binding in Sulfobetaine Micelles. Effects on Acid Equilibria and Rates Daniel W. Tondo,† Jonas M. Priebe,† Bruno S. Souza,† Jacks P. Priebe,† Clifford A. Bunton,‡ and Faruk Nome*,† Departamento de Quı´mica, UniVersidade Federal de Santa Catarina, Floriano´ polis-SC 88040-900, Brazil, and Department of Chemistry and Biochemistry, UniVersity of California, Santa Barbara, California ReceiVed: July 4, 2007; In Final Form: September 10, 2007
The rate of specific hydrogen ion catalyzed hydrolysis of 2-(p-heptoxyphenyl)-1,3-dioxolane and equilibrium protonation of 1-hydroxy-2-naphthoate ion in zwitterionic micelles of SB3-14, C14H29NMe2+(CH2)3SO3-, are increased markedly by NaClO4 which induces anionic character and uptake of H3O+ in the micelles. Other salts, for example, NaNO3, NaBr, and NaCl, have similar but much smaller effects on this uptake of H3O+.
The observation that anionic micelles increase the equilibrium protonation of bases by H3O+, and that cationic micelles decrease it, was explained by Hartley in terms of charge driven interactions between ions and ionic micelles, and nonionic micelles apparently had small effects on acid-base equilibria.1 These generalizations, sometimes called Hartley’s rules, applied to the kinetics of bimolecular reactions involving ionic reagents, and quantitative treatments of transfer equilibria between water and micelles as discrete reaction regions, with ion-specific competition between reactive and inert counterions, were applied to reaction rates and equilibria in solutions of micelles and other association colloids.1 Zwitterionic sulfobetaine micelles, for example, have no overall charge, but there is a charge gradient between the interior and the interfacial surface, and betaine micelles interact with anions.2-4 The interaction is ion-specific; it is weak for high charge density, strongly hydrated ions, for example, OH-, and increases with decreasing ionic charge density, following the Hofmeister series and the Pearson hard-soft classification. This anion binding has been treated theoretically in terms of electrostatic and ion-specific interactions.5 There is considerable evidence on the effects of sulfobetaine micelles on spontaneous reactions of anionic substrates and on bimolecular reactions with anionic reagents, but as far as we know, there is no kinetic evidence on hydrogen ion catalyzed reactions, although they are well studied in ionic micelles.1 Perchlorate ion binds effectively to sulfobetaine micelles, and the anion order is ClO4- > NO3- ≈ Br- > Cl- > OH-, as shown by the use of selective ion electrodes (ISE), electrophoresis, ionic conductance, NMR spectroscopy, and kinetic evidence.5b,c,6 Anion affinities follow Langmuir isotherms, but the extent of anionic coverage, θmax, is much lower than that with cationic micelles, and for NaClO4 in N-tetradecyl-N,Ndimethylammonio-1-propanesulfonate (SB3-14), it is ca. 0.2.7,8 Extents of anionic coverage from zeta potentials, ζ, and ISE follow the above anion order, and with added NaClO4, the † ‡
Universidade Federal de Santa Catarina. University of California, Santa Barbara.
10.1021/jp075208m CCC: $37.00
SCHEME 1
SCHEME 2
sulfobetaine micelle becomes modestly anionic, but is only weakly anionic on addition of NaCl or NaBr.7 Anionic micelles effectively incorporate cations, for example, H3O+, in competition with other cations, and with minor anion effects.1 The situation is different for sulfobetaine micelles, which should have little affinity for cations, for example, Na+ or H3O+, that do not bind specifically. However, sulfobetaine micelles, in the presence of ClO4-, become anionoid and should bind H3O+, but Cl-, for example, should have little effect.5-8 We have tested this hypothesis by examining both the equilibrium dissociation of 1-hydroxy-2-naphthoic acid (HNA), Scheme 1, and the kinetics of the acid hydrolysis of 2-(pheptoxyphenyl)-1,3-dioxolane (HPD), in sulfobetaine micelles with very dilute HCl or HClO4, ca. 10-3 to 0.1 M, Scheme 2. The dioxolane was selected because the acid hydrolysis is first order with respect to substrate concentration, the mechanism and acid catalysis are well understood, and the reaction is accelerated by anionic surfactants.9,10 Added electrolytes increase the protonating power of moderately concentrated acids (>1 M), beyond the range of the pH scale, where Hammett’s acidity scale is applicable, but they should have little effect in aqueous media at pH > 1. The hydrophobic indicator and the dioxolane (concentrations of 8.0 × 10-5 and 1.1 × 10-5 M, respectively), are sparingly soluble in water and should be almost completely micellar-bound. For the hydrolysis and indicator experiments in 0.05 M SB3-14, all at 25.0 ( 0.1 °C, acid hydrolysis was followed kinetically by the increasing absorbance at 286 nm and equilibrium protonation by the increasing absorbance at 358 nm. © 2007 American Chemical Society
11868 J. Phys. Chem. B, Vol. 111, No. 41, 2007
Figure 1. Plots of log kobsd for the hydrolysis of 2-(p-heptoxyphenyl)1,3-dioxolane (HPD) against log[HCl] (9) and log[HClO4] (b), 0.05 M SB3-14, at 25.0 °C.
Figure 2. Salt effects on log krel for the hydrolysis of HPD with 0.005 M HCl and 0.05 M SB3-14 at 25.0 °C: NaCl (9); NaBr (∆); NaNO3 (1); NaClO4 (b). Without added salt, kobsd ) 2.5 × 10-3 s-1.
Figure 1 shows increases in the first-order rate constants, kobsd, with added HClO4 and HCl for the specific hydrogen ion catalyzed, A1, hydrolysis of 1.1 × 10-5 M HPD, in the presence of SB3-14. The reaction with HClO4 is considerably faster than that with HCl, and therefore, even very dilute ClO4- is more effective than Cl- in increasing the incorporation of H3O+ into the micelles. From the plot of kobsd versus [HX] for HCl, the second-order rate constant, kH+, is 1.18 ( 0.03 M-1 s-1, and for HClO4, it is 10.5 ( 0.96 M-1 s-1, over a range of [H+] (Figure 1). At constant [HCl] (0.005 M), salts markedly increase the rate constant, krel, relative to that without salt (Figure 2), and the effect follows the Hofmeister series. Comparison of the results in Figures 1 and 2 shows that our general conclusions are not affected by the presence of dilute Na+. There is the possibility that the salts are affecting acidity in the aqueous pseudophase, and specific kinetic effects are observed with electrolytes in moderately concentrated strong acids.11 However, these effects are seen only at salt and acid concentrations very much higher than those used in the present work, and we eliminate their possible role by examining equilibrium protonations with dilute electrolyte.11 In the micellar solution, where the anionic coverage,8 θmax, for NaClO4 in SB3-14 is ca. 0.2, incorporation of salts will increase the local ionic strength in the micellar pseudophase by less than 20% and much less than observed effects on both acidity and kinetics. Evidence on the micellar-mediated dissociation equilibria of HNA is in Figure 3, with the absorbance at 358 nm increasing
Letters
Figure 3. Plot of the absorbance of HNA (8.0 × 10-5 M) at 358 nm vs pH (O) controlled by addition of HCl and vs concentration of NaCl (9), NaBr (2), NaNO3 (1), and NaClO4 (b), in the presence of 0.001 M HCl, at 25.0 °C.
Figure 4. Plot of the apparent pH against total salt concentration, with 0.001 M HCl and [SB3-14] ) 0.05 M, at 25.0 °C: NaCl (9); NaBr (∆); NaNO3 (1); NaClO4 (b).
on protonation of the anionic form of the indicator (NA-). Spectrophotometric titration of fully micellized HNA is shown in Figure 3, where the open points show formation of NA- as a function of pH, in 0.05 M SB3-14 and no added salt, and the data give an apparent dissociation constant, pKaapp, for HNA of 2.61 ( 0.03. This calculated value is slightly lower than that of 2.88 ( 0.02 at 25.0 °C calculated in aqueous solution from a similar spectrophotometric titration (data not shown). The solid points represent absorbance as a function of added salt for a solution of the indicator and 0.001 M HCl with 0.05 M SB314. The salt order on the equilibrium dissociation of HNA is similar to that observed in acid catalyzed dioxolane hydrolysis. The apparent pH in the interfacial region of the SB3-14 micelles can be estimated from the data in Figure 3 by applying equation 1, which has been applied to pHapp measurements in micelles,12,13 and the results are shown in Figure 4. The increase in the apparent local hydrogen ion concentrations in going from NaCl to NaClO4 is consistent with the kinetic data in Figures 1 and 2, showing that specific anion binding markedly increases hydrogen ion concentration in the micellar interfacial region.
pHapp ) pKaapp + log([NA-]/[NAH])
(1)
As shown in cartoon form in Scheme 3, the effects of NaClO4, relative to NaCl, or other salts, are due to development of negative charge in the anion-modified interfacial region of sulfobetaine micelles which strongly favors an increase in hydrogen ion concentration and therefore affects both indicator
Letters
J. Phys. Chem. B, Vol. 111, No. 41, 2007 11869
SCHEME 3: Chameleon-like Behavior of Zwitterionic Micelles
protonation and rates of A1 hydrolysis in that region. Although the mechanism of the Hofmeister effect remains to be elucidated, our results are consistent with the observed ion-induced changes in packing of monolayers,14 and as a result of incorporation of ClO4-, the sulfobetaine micelle behaves as an anionic micelle, such as that of sodium dodecyl sulfate (SDS), which has a higher fractional micellar charge, and where indicator protonation and acid hydrolyses are well studied.1,9,10 It seems from previous evidence14 and the hereby reported results that the incorporation of anions in zwitterionic surfaces may well be important to understand the behavior of biological membranes. Indeed, added anions promote changes in both packing and surface charges which may in turn modulate the properties of the zwitterionic surface. Micellar effects upon rates of A1 hydrolyses depend on local concentrations of H3O+ and the substrate, and the rate constant, in the micellar pseudophase, but protonation of NA- is an equilibrium process, and it is reasonable to relate both effects to anion-induced increased acidity in the micellar pseudophase. Under some conditions, betaine micelles incorporate anions, but not cations, although incorporation of anions such as perchlorate promotes the uptake of H3O+ into the micelles, and we can control this chameleon-like behavior by choice of the added salt. Experimental Materials. We used N-tetradecyl-N,N-dimethylammonio-1propanesulfonate (SB3-14) (Sigma) and 1-hydroxy-2-naphthoic acid (NAH) (Aldrich). Preparation and purification of of 2-(pheptoxyphenyl)-1,3-dioxolane (HPD) are described.9,10 Solutions were prepared immediately before use. All measurements were made at 25 °C. Other reagents and solvents were of analytical grade and were used without further purification. Kinetics and Spectroscopic Measurements. All measurements were made with diode-array spectrophotometers, with thermostated cell holders, in aqueous sulfobetaine, SB3-14, under conditions such that the spectroscopic probe NAH and the organic substrate HPD are almost wholly micellar-bound and [surfactant] is much higher than the critical micellar concentration. All pH measurements were made with a Metrohm model 713 pH meter. Reactions were started by adding 30 µL of a stock solution of the substrate (1.1 × 10-3 M) in water to 3 mL of reaction
mixture. Observed first-order rate constants, kobsd, were calculated by nonlinear least-squares fitting of the absorbance-versustime curve, and correlation coefficients were always better than 0.999. Second-order rate constants were from linear plots of kobsd against [acid], and simple programs were used in fitting the nonlinear indicator data. The higher absorbance of aromatic acids and esters relative to the anions is a general phenomenon. Acknowledgment. We are grateful to Capes, PRONEX, CNPq, FAPESC, Brazil, and the National Science Foundation, CHE 0411990, for support of this work. References and Notes (1) Bunton, C. A.; Nome, F.; Quina, F. H.; Romsted, L. S. Acc. Chem. Res. 1991, 24, 357-364. (2) (a) Pillersdorf, A.; Katzhendler, J. Isr. J. Chem. 1979, 18, 330338. (b) Bunton, C. A.; Mhala, M. M.; Moffatt, J. R. J. Phys. Chem. 1989, 93, 854-858. (c) Lee, B.; Nome, F. Langmuir 2000, 16, 10131-10136. (3) Chevalier, J.; Kamenka, N.; Chorro, M.; Zana, R. Langmuir 1996, 12, 3225-3232. (4) (a) Weers, J.; Rathman. J.; Axe, F.; Crichlow, C.; Foland, L.; Scheuing, D.; Wiersema, R.; Zielske, A. Langmuir 1991, 7, 854-867. (b) Savelli, G.; Germani, R.; Brinchi, L. In Reactions and Synthesis in Surfactants Systems; Texter, J., Ed.; Marcel Dekker: New York, 2001; Chapter 8. (5) (a) Baptista, M. S.; Cuccovia, I.; Chaimovich, H.; Politi, M. J.; Reed, W. F. J. Phys. Chem. 1992, 96, 6442-6449. (b) Iso, K.; Okada, T. Langmuir 2000, 16, 9199-9204. (c) Masudo, T.; Okada, T. Phys. Chem. Chem. Phys. 1999, 1, 3577-3582. (d) Ohshima, H. J. Colloid Interface Sci. 1994, 168, 269-271. (e) Yokohama, T.; Macka, M.; Haddad, P. R. Fresenius J. Anal. Chem. 2001, 371, 502-506. (f) Yokohama, T.; Macka, M.; Haddad, P. R. Anal. Chim. Acta 2001, 442, 221-230. (6) (a) Bertoncini, C.; Nome, F.; Cerichelli, G.; Bunton, C. A. J. Phys. Chem. 1990, 94, 5875-5878. (b) Cerichelli, G.; Chiarini, M.; Di, Profio, P.; Germani, R.; Savelli, G.; Mancini, G.; Bunton, C. A.; Gillitt, N. D. Langmuir 1998, 14, 2662-2669. (7) Marte, L.; Beber, R. C.; Farrukh, M. A.; Micke, G. A.; Costa, A. C. O.; Gillitt, N. D.; Bunton, C. A.; Profio, P. D.; Savelli, G.; Nome, F. J. Phys. Chem. B 2007, 111, 9762-9769. (8) Beber, R. C.; Bunton, C. A.; Savelli, G.; Nome, F. Prog. Colloid Polym. Sci. 2004, 128, 249-254. (9) Bunton, C. A.; Wolfe, B. J. Am. Chem. Soc. 1973, 95, 3742-3749. (10) Gonsalves, M.; Probst, S.; Rezende, M. C.; Nome, F.; Zucco, C.; Zanette, D. J. Phys. Chem. 1985, 89, 1127-1130. (11) Paul, M. A.; Long, F. A. Chem. ReV. 1957, 57, 1. (12) Frescura, V. L. A.; Marconi, D. M. O.; Zanette, D.; Nome, F.; Blasko, A.; Bunton, C. A. J. Phys. Chem. 1995, 99, 11494-11500. (13) Zanette, D.; Leite, M. R.; Reed, W.; Nome, F. J. Phys. Chem. 1987, 91, 2100-2102. (14) Gurau, M. C.; Lim, S. M.; Castellana, E. T.; Albertorio, F.; Kataoka, S.; Cremer, P. S. J. Am. Chem. Soc. 2004, 126, 10522-10523.
