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Effects of micellar head group structure on the spontaneous hydrolysis of methyl naphthalene-2-sulfonate. The role of perchlorate ion |
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Journal of the Chemical Society, Perkin Transactions 2,
Volume 1,
Issue 2,
1998,
Page 361-364
Lucia Brinchi,
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摘要:
EVects of micellar head group structure on the spontaneous hydrolysis of methyl naphthalene-2-sulfonate. The role of perchlorate ion Lucia Brinchi,a Pietro Di Profio,a Raimondo Germani,a Gianfranco Savelli,a Nicoletta Spreti b and CliVord A. Bunton c a b c Dipartimento di Chimica Università di Perugia 06100 Perugia Italy Dipartimento di Chimica Ingegneria Chimica e Materiali Università di L’Aquila 67010 L’Aquila Italy Department of Chemistry University of California Santa Barbara CA 93106 USA The spontaneous (SN2) hydrolysis of methyl naphthalene-2-sulfonate (MeONs) in water is inhibited by cationic anionic and zwitterionic micelles of the following surfactants CTAOMs n-C16H33N1Me3MeSO3 2; CTPAOMs n-C16H33N1Pr3MeSO3 2; SDS C12H25OSO3 2Na1; SB3-14 n-C14H29N1Me2(CH2)3SO3 2;- SBBu3-14 n-C14H29N1Bu2(CH2)3SO3 2; DMMAO n-C14H29N1Me2O2; DPMAO n-C14H29N1Pr2O2.Rate constants krel relative to those in water are in the range 0.55–0.63 for all the cationic and zwitterionic micelles including the protonated amine oxides. The value of krel in anionic micelles of SDS is 0.22 but NaClO4 sharply decreases krel in SB3-14 from 0.56 to 0.15. These rate effects are not related to variations in substrate binding but depend upon interactions of the head groups with the initial and transition states. Introduction Aqueous micelles affect rates of nonsolvolytic bimolecular reactions by controlling concentrations of the two reagents at the micelle–water interface. Therefore overall rate constants depend upon transfer equilibria between water and micelles and second-order rate constants in each region.1 Nucleophilic reactions have been widely studied and second-order rate constants at micelle–water interfaces are typically similar to or slightly lower than those in water.For SN2 reactions of Br2 or Cl2 mediated by cationic micelles of quaternary ammonium ion surfactants rate constants at the micelle–water interface increase modestly with increasing head-group bulk of the surfactant, 2 but the effect is inverted for reactions of OH2.3 Rates of intramolecular SN2-like cyclisations at micellar surfaces also increase modestly with increasing head group size.4 Spontaneous bimolecular hydrolyses of alkyl halides and sulfonate esters are mechanistically simple because the solvent e.g.water participates only nucleophilically and by solvating leaving groups and not as a general acid–base catalyst as in some deacylations. In the transition state positive charge develops on the nucleophilic water molecule and is dispersed into the solvent by hydrogen-bonding. However despite involvement of water molecules as nucleophiles or general bases aqueous micelles do not strongly affect rates of either deacylations or SN2 hydrolyses indicating that the micellar interfacial region is ‘water-rich’.1,5 There could also be effects due to the high ionic concentration in this region as well as its slightly lower polarity relative to water.1,6,7 This region is electrically asymmetric in ionic and zwitterionic micelles which may affect hydrolyses in which local charges are developed in the transition state and inhibition of spontaneous hydrolyses is larger in anionic than in cationic micelles.5c,7 Inhibition is larger for SN1 than for spontaneous bimolecular hydrolyses and then the more effective inhibitors are cationic rather than anionic micelles.5a,d,8 These different charge effects can be rationalized in terms of interactions of the micellar head groups with local charges in the transition states of uni- and bi-molecular hydrolyses at acyl or alkyl centers.Similar interactions with head groups are apparently important in anionic decarboxylations 9 and dephosphorylations 10 and E1cB reactions 11 in ionic and zwitterionic micelles. Although charge affects rates of spontaneous hydrolyses at micellar surfaces there is limited evidence regarding effects of head group size or structure.10 Engberts and co-workers found that micelles and other association colloids modestly inhibit spontaneous hydrolyses of acyltriazoles and the extent of inhibition is sensitive to colloid structure.5b,c Spontaneous hydrolyses of benzenesulfonyl chlorides are micellar inhibited and inhibition is greater for anionic than for cationic micelles but zwitterionic sulfobetaine and cationic micelles behave similarly.12 Similar micellar charge effects are seen in spontaneous hydrolyses of carboxylic anhydrides and acyl chlorides.5d,8a In the present work we examine micellar effects on the spontaneous (SN2) hydrolysis of methyl naphthalene-2-sulfonate (MeONs) Scheme 1.SO3Me + H2O Scheme 1 This reaction is a mechanistically simple concerted displacement whereas hydrolyses of acyl derivatives may be concerted or step-wise,13 and the hydrophobic substrate binds strongly to micelles. As a result we obtain limiting values of the rate constants in relatively dilute surfactant so that changes due to micellar growth are unimportant. We used cationic anionic and zwitterionic surfactants and varied the head group structures. The amine oxides are useful because protonation converts them from zwitterions which are nucleophilic,14 into cations and protonation in dilute acid is well established.15 Surfactants used in this work are cetyltrimethylammonium methanesulfonate (CTAOMs) n-C16H33N1Me3MeSO3 2; cetyltripropylammonium methanesulfonate (CTPAOMs) n-C16H33N1Pr3MeSO3 2; sodium dodecyl sulfate (SDS) C12H25OSO3Na; 3-(N,N-dimethylmyristylammonio) propanesulfonate (SB3-14) n-C14H29N1- Me2(CH2)3SO3 2; 3-(N,N-dibutylmyristylammonio)propanesulfonate (SBBu3-14) n-C14H29N1Bu2(CH2)3SO3 2; 1,1-dimethylmyristylamine oxide (DMMAO) n-C14H29N1Me2O2; 1,1- dipropylmyristylamine oxide (DPMAO) n-C † Cetyl = hexadecyl myristyl = tetradecyl.J. Chem. Soc. Perkin Trans. 2 1998 SO3 – + H+ + MeOH 14H29N1Pr2O2.† 361 Table 1 Rate constants of the spontaneous hydrolysis of MeONs in cationic surfactants CTAOMs CTPAOMs DMMAO [Surfactant]/1023 mol dm23 [Surfactant]/1023 mol dm23 k k obs/1026 s21 obs/1026 s21 12.3 8.04 7.58 7.27 12.3 9.30 8.83 7.82 7.35 7.34 0 3 5 10 30 50 6.90 6.96 7.32 100 Table 2 Rate constants of the spontaneous hydrolysis of MeONs in sulfobetaine surfactants [Surfactant]/1023 mol dm23 SB3-14 SBBu3-14 Table 4 Rate constants of the spontaneous hydrolysis of MeONs in SDS k k obs/1026 s21 obs/1026 s21 12.3 12.3 9.15 8.12 0 1 3 5 8.50 8.07 7.69 7.70 7.75 7.19 7.09 6.89 7.17 7.11 8 10 30 50 100 The counterion in the cationic surfactants was mesylate rather than halide ion to limit nucleophilic participation,2 as in solutions of HBr and DMMAO or DPMAO,16 and therefore we used MeSO3H to protonate the amine oxides.(1) 4 Results and discussion Micellar kinetics Micelles and water are treated as distinct reaction media i.e.as pseudophases,1 and the observed first-order rate constant kobs is given by eqn. (1) kobs = k9W 1 k9MKS[Dn] 1 1 KS[Dn] Substrate S is rapidly partitioned between water and micelles designated by subscripts W and M respectively with a binding constant KS with respect to micellized surfactant (detergent) Dn whose concentration is the total less that of the monomer i.e. the critical micelle concentration c.m.c. under the kinetic conditions.17 The value of KS of MeONs in cationic micelles2,3 is ca. 103 dm3 mol21 and this hydrophobic substrate probably promotes micellization so that except in dilute surfactant MeONs is fully micellar bound and k9M ª kobs.We observed limiting rate constants in the range 0.03–0.1 mol dm23 surfactant consistent with quantitative substrate binding (Tables 1–4). The electrolyte concentration was varied by addition of MeSO3Na or MeSO3H and NaClO4 with SB3-14 and the hydrolysis rate in water is almost unaffected by up to 1 mol dm23 electrolyte. Rate data with added MeSO3Na and MeSO3H are available as supplementary material and in Fig. 1.‡ Added C1O2 binds very strongly to sulfobetaine micelles as shown by 35Cl NMR spectroscopy and strong inhibition of the reaction of Br2 with ‡ This material has been deposited in the Supplementary Publications Scheme (SUPPL NO. 57320 3 pp.) For details of the deposition scheme see ‘Instructions for Authors’ J. Chem. Soc. Perkin Trans.2 available via the RSC Web pages (http://www.rsc.org/authors). J. Chem. Soc. Perkin Trans. 2 1998 362 Table 3 Rate constants of the spontaneous hydrolysis of MeONs in amine oxide surfactants DPMAO k kobs/1026 s21 a kobs/1026 s21 b obs/1026 s21 a 13.3 9.03 13.4 9.74 13.4 9.08 7.42 7.61 7.20 7.97 7.64 0 2 6 7 10 12 7.12 7.04 7.07 6.90 6.96 7.06 18 21 31 35 6.72 7.47 a In 0.05 mol dm23 MeSO3H. b In 0.1 mol dm23 MeSO3H. [SDS]/1023 mol dm23 kobs/1026 s21 12.5 7.43 3.54 3.10 2.85 2.93 0 10 20 30 50 70 2.69 100 Fig. 1 Rate constants of the spontaneous hydrolysis of MeONs in 0.05 mol dm23 SB3-14 with addition of MeSO3Na (d) and NaClO4 (j) MeONs.16 It therefore converts a zwitterionic into an anionic micelle which should affect rates of spontaneous hydrolyses.Values of k9M are given in Table 5 and variations of kobs with M were [surfactant] are given in Tables 1–4. Values of k9 obtained by fitting rate data to eqn. (1) as described.1,2 EVect of micellar charge The spontaneous hydrolysis of methyl benzenesulfonate is inhibited by both CTAOMs and SDS and values relative to reaction in water krel are 0.70 and 0.40 respectively.5d,18 Values of krel for reaction of MeONs in surfactants (Table 5) are similar to but slightly lower than those for reaction of methyl benzenesulfonate despite differences in hydrophobicities of these substrates (for the benzenesulfonate 18 KS = 50–70 dm3 mol21).Both substrates reside in a similar region of the micelle Table 5 Rate constants of the spontaneous hydrolyses in the micellar pseudophase a Surfactant b k9M/1026 s21 4 CTAOMs CTAOMs 1 MeSO3Nab CTPAOMs CTPAOMs 1 MeSO3Nab SB3-14 SB3-14 1 MeSO3Nab SB3-14 1 MeSO3Hb SB3-14 1 NaClO c SBBu3-14 SBBu3-14 1 MeSO3Nab DMMAOH1 1 MeSO3Hd DMMAOH1 1 MeSO3Hb DPMAOH1 1 MeSO3Hd SDS 7.33 (0.59) 7.23 (0.58) 6.93 (0.55) 7.01 (0.56) 7.03 (0.56) 7.94 (0.64) 7.27 (0.58) 1.85 (0.15) 7.73 (0.62) 7.40 (0.59) 7.22 (0.58) 6.81 (0.54) 7.06 (0.56) 2.81 (0.22) a At 25.0 8C with k9W = 1.25 × 1025 s21. Values in parentheses are relative rate constants in micelles and water krel. b 0.1 mol dm23 MeSO3Na or MeSO3H.c 1–1.6 mol dm23 NaClO4 (Fig. 1). d 0.05 mol dm23 MeSO3H. which based on NMR data with MeONs is the micelle–water interface adjacent to the ionic head groups.19 However increased penetration of the more hydrophobic substrate MeONs should partially shield it from water molecules.19 Zwitterionic and cationic micelles including protonated amine oxides behave differently from anionic micelles (Table 5). Neutral amine oxides can increase kobs by nucleophilic dis- M in Table 5 do not seem to be related to placement on MeONs14 and we discuss these reactions elsewhere. The values of k9 the hydrogen bonding of water to anionic or zwitterionic head groups. Hydrogen bond donation to sulfate or sulfonate residues might activate water as a nucleophile,20 but this effect does not explain our results because reactions are slower in anionic than in cationic micelles and water at the micellar surface is apparently not deactivated by hydrogen-bonding to protonated amine oxides (Table 5).However analyses of micelle–solute interactions in terms of linear solvation free energy relationships indicate the importance of hydrogen bonding,7 and the specificity of micelle–ion interactions 1 is related to the ease of partial ionic dehydration.21 EVect of head group structure The insensitivity of k9M to head group bulk and structure (Table 5) was unexpected because of the observed acceleration in spontaneous anionic decarboxylations and dephosphorylations in cationic and sulfobetaine micelles with increasing head group bulk.9b,10 Cyclization of the phenolic derivatives 1 (Scheme 2) is also O– O (CH2)3 + X– O O(CH2)3X 1a X = Br b X = I Scheme 2 accelerated by an increase in head group bulk,4 as are SN2 reactions of Cl2 and Br2 with MeONs.2 However an increase in the bulk of cationic head groups slightly decreases second-order rate constants for reactions of OH2 at micellar surfaces.3 Bulky alkyl groups at the ammonium ion center may partially exclude water molecules from the micellar surface and decrease its polarity and hydrogen bonding of interfacial water to anions.The absence of such an acceleration of reactions of OH2 is understandable because this ion has a high affinity for water and its hydration should be unaffected by an increase in head group bulk.The Hughes–Ingold rules predict that SN2 solvolyses should be inhibited by a decrease in water content or polarity of the solvent.22 Extension of these rules to micellar surfaces is satisfactory for a number of reactions 1d,5 and we therefore expected that polarity and availability of water at micellar surfaces and therefore rate constants would be affected by bulky alkyl groups. Similarities in kinetic behaviours of quaternary ammonium and sulfobetaine micelles are understandable. The charge in ionic micelles is extensively neutralized by counterions 1 and if the trimethylene tether in our sulfobetaine surfactants is extended 23 the sulfonate residue should not strongly perturb the environment of the quaternary ammonium ion.As noted earlier protonated amine oxides are cationic surfactants but although they can strongly hydrogen-bond to water they appear to have no special effect on its reactivity towards MeONs (Table 5). H1 R9R2N1]OH R9R2N1]O2 The nature of the reaction region Rate enhancements of nonsolvolytic bimolecular reactions at micellar surfaces are fitted quantitatively by models that take into account the high concentration of reagents in the interfacial region which is often identified as the Stern layer a few Angstroms thick where counterions are concentrated.1 Based on other kinetic evidence the properties of this region as a reaction medium should depend on the head group structure cf. refs. 2–5 8–10 in contrast to the apparent insensitivity of our spontaneous SN2 hydrolyses (Table 5).The hydrophobic naphthalene group of MeONs will orient itself towards the quaternary ammonium ions and the attached alkyl tails,19,24 but the methylsulfonate residue can extend into a more aqueous region (the question of the ‘wetness’ of the micellar surface and the extent of water–hydrocarbon contact has been discussed in terms of various micellar models 1,5a,7,25). In the transition state negative charge should build up on the naphthalenesulfonate residue and interact unfavorably with anionic head groups in SDS micelles consistent with the difference in krel for these and the cationic or zwitterionic micelles. (Table 1 and refs. 5d 8 11). Inhibitions are similar with cationic and zwitterionic micelles which have similar charge asymmetry in the interfacial region.There is no indication that changes in the average location of a substrate in the micellar interfacial region as related to its hydrophobicity have a major effect on the rate of spontaneous hydrolysis. 4 4 4 4 The role of perchlorate ion The hydrolysis rate of MeONs in SB3-14 is almost unaffected by the addition of up to 1.5 mol dm23 MeSO3Na (with a slight increase) but we observe a large effect upon addition of NaClO4 (Fig. 1). As noted perchlorate ion interacts strongly with sulfobetaine micelles,16 and in effect C1O2 converts a zwitterionic sulfobetaine into an anionic micelle. This generation of head group charge inhibits the spontaneous hydrolysis of MeONs (Table 5 and Fig.1). However the inhibition by NaClO4 is higher than expected for an anionic micelle with krel = 0.15 as compared with 0.22 in SDS. It is difficult to explain inhibition by C1O2 solely in terms of micellar charge because micellized SB3-14 1 C1O2 should not be more ‘anionic’ than SDS. Chevalier and co-workers have suggested that the interfacial region of betaine micelles can be very open,23,26 which allows water molecules to penetrate these regions. However C1O2 interacts with the ammonium centers based on changes in the 14N NMR spectrum,16 and it will then expel water from this region and thus inhibit spontaneous bimolecular hydrolyses (Table 5). The insolubility of cationic surfactants 1 C1O2 prevents our examining this system.4 Conclusions Our evidence regarding effects of cationic betaine sulfonate J. Chem. Soc. Perkin Trans. 2 1998 363 and anionic micelles on a spontaneous SN2 hydrolysis fits the generalization that cationic and betaine micelles behave similarly as reaction media.9b,10,11 Rate enhancements of reactions of anionic reagents are generally lower with betaine sulfonate than with cationic micelles but these differences are due to relatively weak binding of anions to betaine micelles and not to reactivities at micellar surfaces.27 We can explain rate effects on a variety of spontaneous reactions in terms of a lower polarity and water availability in the interfacial region relative to bulk water,6,7,23 together with electrical asymmetry in this region which generates the differences between cationic and anionic micelles.5d The E1cB reaction of the carbanion of a fluorenyl carboxylic ester also illustrates the role of charge asymmetry at micellar surfaces.In this reaction negative charge is dispersed out of the hydrophobic fluorenyl group in the ketene-like transition state. Rate effects are small but reaction is inhibited by cationic and betaine micelles and accelerated by anionic and phospholipidderived micelles.11 We note that micellar effects on rates of spontaneous reactions may be very large as in decarboxylations,9 or small as in these SN2 hydrolyses or the E1cB reaction.11 They are qualitatively understandable in terms of a simple model of the micelle–water interface and evidence on its hydration and polarity.6,7 These considerations should also apply to reactions in other association colloids e.g. microemulsions and vesicles.1,5c However changes in the properties of the interfacial region induced for example by changes in the head group structure can induce modest rate effects depending on reaction mechanism. Experimental Materials Preparation and purification of MeONs and the surfactants have been described.2,10 Critical micelle concentrations (c.m.c./ mmol dm23) of the zwitterionic surfactants were SB3-14 0.29; SBBu3-14 0.11; DMMAO 0.14; DPMAO 0.05. They were measured by surface tension and there were no minima in the relevant plots.1a Reactions were carried out in redistilled deionized water. Kinetics Reactions were followed at 25.0 8C in a Shimadzu UV-160 A or an HP 8452 spectrophotometer by following decreasing absorbance at 326 nm with 1024 mol dm23 MeONs as described.2 The slower reactions could not be followed to 10 half-lives and kobs was then calculated by a nonlinear leastsquares fitting of the variation of absorbance with time to a first-order rate equation.For the faster reactions values of kobs from this method and those based on an infinity absorbance agreed. Values of kobs over a range of [surfactant] are in Tables 1–4 and supplementary material. Addition of MeSO3H to solutions of the sulfobetaines had little effect on kobs showing that there is no reaction with OH2. Reactions in solutions of amine oxides were followed in excess MeSO3H where protonation is quantitative.15 The electrolyte concentration was varied in the absence of surfactants by addition of MeSO3H MeSO3Na and NaClO4; the hydrolysis rate in water is almost unaffected by added salt.In 0.05–0.1 mol dm23 MeSO3H kobs = 1.33 ± 0.01 × 1025 s21; in 0.1–1 mol dm23 MeSO3Na kobs = 1.25 ± 0.07 × 1025 s21; in NaClO4 0.1–1 mol dm23 kobs = 1.14 ± 0.40 × 1025 s21. Acknowledgements Support of this work by Consiglio Nazionale delle Ricerche Rome the Ministero dell’Università e Ricerca Scientifica e Tecnologica Rome and the U.S. Army Office of Research is gratefully acknowledged. 364 J. Chem. Soc. Perkin Trans. 2 1998 Paper 7/05316I References 1 (a) J. H. Fendler Membrane Mimetic Chemistry Wiley-Interscience New York 1982; (b) L.S. Romsted in Surfactants in Solution ed. K. L. Mittal and B. Lindman Plenum Press New York 1984 vol. 2 p. 1015; (c) C. A. Bunton and G. Savelli Adv. Phys. Org. Chem. 1986 22 213; (d ) C. A. Bunton F. Nome F. H. Quina and L. S. Romsted Acc. Chem. Res. 1991 24 357. 2 (a) R. Bacaloglu C. A. Bunton and F. Ortega J. Phys. Chem. 1989 93 1497; (b) R. Bacaloglu C. A. Bunton G. Cerichelli and F. Ortega J. Phys. Chem. 1990 94 5068. 3 C. Bonan R. Germani P. P. Ponti G. Savelli G. Cerichelli R. Bacaloglu and C. A. Bunton J. Phys. Chem. 1990 94 5331. 4 G. Cerichelli L. Luchetti G. Mancini M. N. Muzzioli R. Germani P. P. Ponti N. Spreti G. Savelli and C. A. Bunton J. Chem. Soc. Perkin Trans. 2 1989 1081. 5 (a) F. M. Menger Acc. Chem.Res. 1979 12 111; (b) N. Fadnavis and J. B. F. N. Engberts J. Org. Chem. 1982 47 152; (c) W. H. Noordman W. Blokzijl J. B. F. N. Engberts and M. J. Blandamer J. Chem. Soc. Perkin Trans. 2 1995 1411; (d ) C. A. Bunton in Nucleophilicity ed. J. M. Harris and S. P. McManus Adv. Chem. Ser. No. 215 American Chemical Society Washington DC. 1987 ch. 29. 6 (a) K. A. Zachariasse N. Y. Phuc and B. Kozankiewicz J. Phys. Chem. 1981 85 2672; (b) C. Ramachandran R. A. Pyter and P. Mukerjee J. Phys. Chem. 1982 86 3198; (c) E. J. R. Sudhölter G. B. van der Langkruis and J. B. F. N. Engberts Recl. Trav. Chim. Pays-Bas Belg. 1979 99 73; (d ) R. Zana in Surfactant Solutions. New Methods of Investigation ed. R. Zana M. Dekker Inc. New York 1987 p. 272. 7 (a) M.H. Abraham H. C. Chadha J. P. Dixon C. Rafolos and C. Treiner J. Chem. Soc. Perkin Trans. 2 1995 887; (b) F. H. Quina E. O. Alonso and J. P. S. Farah J. Phys. Chem. 1995 99 11 708; (c) M. H. Abraham H. C. Chadha J. P. Dixon C. Rafolos and C. Treiner J. Chem. Soc. Perkin Trans. 2 1997 19. 8 (a) H. Al-Lohedan C. A. Bunton and M. M. Mhala J. Am. Chem. Soc. 1982 104 6654; (b) C. A. Bunton and S. Ljunggren J. Chem. Soc. Perkin Trans. 2 1984 355. 9 (a) C. A. Bunton M. J. Minch J. Hidalgo and L. Sepulveda J. Am. Chem. Soc. 1973 95 3262; (b) P. Di Profio R. Germani G. Savelli G. Cerichelli N. Spreti and C. A. Bunton J. Chem. Soc. Perkin Trans. 2 1996 1505. 10 F. Del Rosso A. Bartoletti P. Di Profio R. Germani G. Savelli A. Blaskó and C. A. Bunton J.Chem. Soc. Perkin Trans. 2 1995 673. 11 V. R. Correia I. M. Cuccovia M. Stelmo and H. Chaimovich J. Am. Chem. Soc. 1992 114 2144. 12 C. A. Bunton M. M. Mhala and J. R. Moffatt J. Org. Chem. 1985 50 4921. 13 A. Williams Acc. Chem. Res. 1989 22 387. 14 W. P. Jencks Catalysis in Chemistry and Enzymology McGraw-Hill Co. New York 1969 p. 91. 15 J. F. Rathmann and S. D. Christian Langmuir 1990 6 391. 16 P. Di Profio Ph.D. Thesis University of Perugia 1996. 17 F. M. Menger and C. E. Portnoy J. Am. Chem. Soc. 1967 89 4698. 18 H. Al-Lohedan C. A. Bunton and J. R. Moffatt J. Phys. Chem. 1983 87 332. 19 R. Bacaloglu C. A. Bunton G. Cerichelli and F. Ortega J. Phys. Chem. 1989 93 1490. 20 M. C. R. Symons Acc. Chem. Res. 1981 14 179. 21 J. D. Morgan D. H. Napper and G. G. Warr J. Phys. Chem. 1995 99 9458. 22 C. K. Ingold Structure and Mechanism in Organic Chemistry 2nd edn. Cornell University Press Ithaca New York 1969 p. 457 ff. 23 Y. Chevalier and P. Le Perchec J. Phys. Chem. 1990 94 1774. 24 (a) S. J. Bachofer and U. Simonis Langmuir 1996 12 1744; (b) P. J. Kreke L. J. Magid and J. C. Gee Langmuir 1996 12 699. 25 (a) P. Fromherz Ber. Bunsenges. Phys. Chem. 1981 85 891; (b) J. A. Butcher and G. W. Lamb J. Am. Chem. Soc. 1984 106 1217; (c) B. Jonsson P.-G. Nilsson B. Lindman L. Guldbrand and H. Wennerstrom in Surfactants in Solution ed. K. L. Mittal and B. Lindman Plenum Press New York 1984 vol. 1 p. 3. 26 Y. Chevalier N. Kamenka M. Chorro and R. Zana Langmuir 1996 12 3225. 27 (a) M. da Silva Baptista I. Cuccovia H. Chaimovich M. J. Politi and W. F. Reed J. Phys. Chem. 1992 96 6442; (b) C. A. Bunton M. M. Mhala and J. R. Moffatt J. Phys. Chem. 1989 93 854. Received 23rd July 1997 Accepted 25th September 1997
ISSN:1472-779X
DOI:10.1039/a705316i
出版商:RSC
年代:1998
数据来源: RSC
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