J. Pkys. Chem. 1992,96, 1957-1961
1957
Effect of Alcohols on Catalysis by Dodecyl Sulfate Micelles Carlos Bravo, J. Ramdn Leis,* and M. Elena Peiia Departamento de Qujmica Flsica, Facultad de Qujmica, Universidad de Santiago, Santiago de Compostela, Spain (Received: May 6, 1991)
We studied the kinetics of the acid denitrosation of N-nitroso-N-methyl-p-toluenesulfonamide when catalyzed by sodium dodecyl sulfate (SDS) or hydrogen dodecyl sulfate (HDS) micelles in the presence of one of eight alcohols, with equimolar alcohol and surfactant concentrations. In preliminary experiments, the effects of the alcohols on micellar structure were investigated by conductimetric determination of micellar ionization and, in some cases, by fluoroscopicmeasurement of mixed micelle aggregation numbers. Both surfactants always catalyzed the reaction, but whereas catalysis by SDS peaked at a certain concentration of surfactant, catalysis by HDS increased with HDS concentration until reaching a limiting value that was maintained thereafter. The degree. of catalysis decreased as the hydrophobic nature of the alcohol increased. The reaction kinetics were in all cases in keeping with a pseudophase ion exchange model in which the volume of the Stern layer varies explicitly with the quantity of alcohol incorporated in the micelle. Micellesubstrate association constants, bimolecular rate constants in the micellar pseudophase and, for SDS, Na+/H+ exchange constants were calculated. All the kinetic and thermodynamicconstants were independent of which alcohol was used and almost the same as those determined in the absence of alcohol. The change in reaction rate due to the alcohols therefore came about solely because the reagents were diluted in the micellar pseudophase. This finding is discussed in terms of the polarity of the Stern layer in single-component and mixed micelles.
Introduction Considerable attention has been paid in recent years to the influence of alcohols on ionic micellar structures, partly because they are the cosurfactants most commonly employed in the preparation of microemulsions. For this purpose, it is usual to use medium-chain alcohols (propanol to hexanol), which form mixed micelles on distribution between the aqueous and micellar phases.'J The alcohol molecules are inserted between the surfactant molecules, therefore forcing apart the ionic heads of the latter, which reduces the surface charge density of the micelle and hence increases its degree of ionization (a).3 In spite of the vast amount of work-with a great variety of experimental techniques-devoted to the perturbations caused by alcohols upon micellar structure, some questions nevertheless remain unanswered, in particular, whereas the results of experiments with fluorescent probes4 have been interpreted as showing that addition of alcohol displaces water molecules from the micellar palisade layer, so reducing its polarity, electron spin-echo modulation experimentsS have suggested the opposite conclusion, i.e. that alcohol increases the penetration of water into the micellar pseudophase. Addition of alcohol also has kinetic effects; it reduces catalysis apparently due largely to the alcohol by cationic incorporating into the micellar phase and so diluting the reagents inside the micelle. This increase in volume is similar to that observed in the formation of microemulsions, which are in fact often called swollen micelle^.^ In this article, we describe a systematic study of the influence of a variety of alcohols on a well-characterized reaction in the presence of anionic micelles (whose kinetic effects have received (1) Kayase, K.; Hayano, S. Bull. Chem. SOC.Jpn. 1977, 50, 83. (2) Manabe, M.; Shirahama, K.; Kode, M. Bull. Chem. SOC.Jpn. 1976, 49, 2904. (3) See, for example: (a) Zana, R.; Yiv, S.; Strazielle, C.; Lianos, P. J. Colloid Interfuce Sei. 1981, 80, 208. (b) Almgrem, M.; Lofroth, J. E. J . Colloid Interfuce Sei. 1981,81,486. (c) Jain, A. K.; Sing, R. P. B. J . Colloid Interfie Sei. 1981, 81, 536. (d) Abu-Hamdiyyah, M.; Rahman, I. A. J. Phys. Chem. 1985,89,2377. (e) Abu-Hamdiyyah, M.; Kumari, K. J. Phys. Chem. 1990, 94, 2518. (4) (a) Lianos, P.; Zana, R. Chem. Phys. Lett. 1980, 76,62. (b) Lianos, P.;Zana, R. Chem. Phys. Lett. 1980, 72, 161. (c) Zana, R.; Yiv, S.; Strazielle, C.; Lianos, P. J. Colloid Interface Sei. 1981, 80, 208. (5) Baglioni, P.; Kevan, L. J. Phys. Chem. 1987, 91, 1516. (6) Bunton, C. A.; Buzzaccarini, F. de J . Phys. Chem. 1982, 86, 5010. (7) Athanassaki, V.; Bunton, C. A.; Buzzaccarini, F. de; J . Phys. Chem. 1982, 86, 5002. (8) Otero, C.; Rodenas, E. J. Phys. Chem. 1986, 90, 5771. (9) Berthod, A. J . Chim. Phys. 1983, 80, 407.
less attention than those of cationic micelles). The fact that reaction kinetics are influenced by the nature of the medium suggests that such studies may throw light on the changes in micellar structure caused by addition of alcohols, especially as regards the polarity of the Stern layer. They may also serve to test whether the existing models for single-component micelles are valid in the presence of additives and to elucidate the influence of the latter on the various factors affecting reactivity within the micelle. The reaction chosen was the acid denitrosation of N-nitrosoN-methyl-p-toluenesulfonamide(MNTS), whose kinetics in the presence of anionic, cationic, and nonionic surfactants we have recently reported.l0 This reaction is well-suited for the study now described because its kinetics comply perfectly with Romsted's pseudophase ion exchange (PIE) model of micellar catalysis;" the intensity of its catalysis by reactive-ion anionic micelles of hydrogen dodecyl sulfate (HDS) increases with surfactant concentration until it reaches a limiting value when all the substrate is associated with micelles. This behavior confirms the constancy of the fraction of neutralized micellar charge B ('1 - CY),an assumption that is basic to Romsted's model and that has been challenged by kinetic and conductimetric evidence of the nonconstancy of B in the case of reactive-ion cationic micelles with OH- or F counterions.12 In this study, we used both reactive-ion micelles of HDS and inert-ion micelles of sodium dodecyl sulfate, SDS. The alcohols used were 1-butanol, 2-butanol, isobutyl alcohol, tert-butyl alcohol, 1-pentanol, tert-amyl alcohol, 1-hexanol, and benzyl alcohol. The alcohol/surfactant mole ratio was always unity when alcohol was present. Prior to the kinetic experiments, we determined the value of B for all the systems studied, since this parameter appears explicitly in the kinetic equations of the PIE model. Experimental Section
HDS was prepared from SDS solutions by ion exchange on a cationic resin, as described elsewhere.I0 HDS solutions were used (10) Bravo, C.; Herva, P.; Leis, J. R.; Peiia, M. E. J . Phys. Chem. 1990, 94, 8816. ( 1 1) (a) Romsted, L. S. In Surfucfantsin Solution; Lindman, B., Mittal, K. L., Eds.; Plenum Press: New York, 1984; Vol 2, p 1015. (b) Romsted, L. S. J. Phys. Chem. Ins,89,5107,5113. (c) Bunton, C. A.; Savelli, G. Adv. Phys. Org. Chem. 1986, 22, 213. (12) (a) Vera, S.;Rodenas, E. Tefmhedron 1986,42, 143. (b) AI-Lohedan, H. A. J. Chem. Soc., Perkin Trans. 2 1989, 1181. (c) Santana Neves, M. de F.; Zanette, D.; Quina, F.; Tadeu Moretti, M.; Nome, F. J . Phys. Chem. 1989, 93, 1502. (d) Rodenas, E.; Vera, S. J . Phys. Chem. 1986, 90, 3414.
0022-3654/92/2096- 1957$03.00/0 0 1992 American Chemical Society
Bravo et al.
1958 The Journal of Physical Chemistry, Vol. 96, No. 4 , 1992 TABLE I: Values of j3 Measured Conductimetricallyfor the Various Systems Studied
system SDS SDS/ 1-butanol SDS/isobutyl alcohol SDS/Z-butanol SDSlrert-butyl alcohol SDS/ 1-pentanol SDSltert-amyl alcohol SDS/ 1-hexanol SDS/benzyl alcohol
P 0.72
0.58 0.57 0.55 0.60 0.55
0.60 0.59 0.59
system HDS/l-butanol HDS/isobutyl alcohol HDS/2-butanol HDSltert-butyl alcohol HDS/ 1-pentanol HDSltert-amyl alcohol HDS/ 1-hexanol HDS/benzyl alcohol
P 0.66 0.62 0.63 0.66 0.63 0.63 0.54 0.63
Figure 2. Results of pyrene quenching experiments. In (fo/Z) plotted against concentrations of N-cetylpyridinium chloride for micellar solutions of ( 0 )HDS-n-hexanol ([HDS] = [n-hexanol]= 0.0992 M); (0) pure SDS ([SDS] = 0.0727 M); (A) HDS-1-hexanol ([HDS] = [lhexanol] = 0.0248 M). TABLE Ik Aggregation Numbers Determined by Fluorescence Ouenchintz Exwriments ~
10' [HDS]/M Figure 1. Dependence of conductivity on HDS concentration in the presence of 1-butanol.
shortly after preparation to avoid any interference from acid hydrolysis. All other reagents were Merck or Aldrich products of the highest commercially available purity and were used as supplied. Kinetics were studied in a Beckman DU-65 spectrophotometer with a thermostated cell carrier by recording the decrease in the 250-nm absorbance due to the consumption of MNTS. Conductivity measurements were performed in a Crison 525 conductimeter using solutions prepared with double-distilled water purified by a Millipore system. All kinetic and conductimetric determinations were carried out a t a temperature of 25.0 f 0.1 OC. Fluorescence measurements were performed with a Perkin-Elmer 560-40 spectrofluorimeter. All kinetic experiments were performed with MNTS concentrations much smaller than that of hydrochloric acid. The absorbance-time data of all kinetic experiments fitted the first-order integrated equation In (A, - A,) = In (A, - A,) - k,t (1) where A,, A,, and A, are the absorbances at times t , 0, and infinity and ko is the first-order pseudoconstant. Values of ko were reproducible to within a f3%.
surfactant SDS (7.23 X M) SDS (6.10 X M) HDS (4.96 X loT2M) M) HDS (2.48 X M) HDS (9.92 X
~
~~~
additive
ii
69
1-hexanol (6.10 X
M)
55
1-hexanol (2.48 X 1-hexanol (9.92 X
M) M)
55 55 58
M). The values of @ obtained for the various mixtures containing alcohol were all less than that determined in the absence of alcohol (Table I), confirming that addition of alcohol increases micellar ionization. The conductivity-[HDS] plots showed two breakpoints (Figure l ) , one at about 7 X M, which we took to be the cmc, and the other, less pronounced change at about (4-5) X lo-* M. We decided to investigate whether the cause of the latter was, as in many other cases of second slope changes in plots of physical quantities against surfactant concentration, a change in micellar structure (hypothetically due to the addition of alcohol, since 5 X lo-* M is too small a concentration for such a change in the absence of additives). To this end, we determined aggregation numbers at various concentrations of alcohols (including zero) and, for the purpose of comparison, carried out a similar study on SDS-alcohol micelles. Following Turro and Yekta,I4 aggregation numbers were determined by measuring the quenching of a micelle-bound fluorescent probe by the binding of a quencher. This technique assumes that the numbers of both probe molecules and quencher molecules per micelle have Poisson distributions, which leads to the expression14
Results and Discussion Detenninstionof 8. The degree of micelle charge neutralization, @, was determined conductimetrically for mixtures with a surfactant/alcohol mole ratio of 1. SDS-alcohol systems had conductivity-[surfactant] profiles consisting of two linear segments meeting at the critical micelle concentration (cmc); the ratio between their slopes was taken as the degree of micellar ionization, a,I3 and @ was calculated as 1 - a. The linearity of the conductivity-[surfactant] curve a t concentrations greater than the cmc meant that under the conditions employed (i.e. with equal mole concentrations of surfactant and alcohol) @ was constant over the whole range of surfactant concentrations studied (0-0.1
where Io and I are the emitted light intensities with quencher concentrations zero and [Q], respectively, ii is the mean surfactant aggregation number, and [D] is the total concentration of surfactant; the mean aggregation number, 2, is calculated from the slope of plots of In (Zo/l)against [Q] for fixed [D]. The probe used was pyrene (at a concentration small enough to prevent excimer formation), and the quencher N-cetylpyridinium chloride.I5 The excitation wavelength was 337 nm, and fluorescence was measured at 394 nm. Figure 2 shows typical In (Io/Z)-[Q] plots. The aggregation number obtained for SDS in the absence of alcohol, 69 (Table 11),agrees well with the value of 60 reported by othersI4 who have used this technique and with the value of 63 obtained in membrane osmometry studies.I6 The values
(13) (a) Hoffmann, H.; Ulbricht, W. Z.Phys. Chem. (Munich) 1977,106, 167. (b) Hoffmann, H.; Tagesson, B. Z. Phys. Chem. (Munich) 1978, 110, 8.
(14) Turro, N. J.; Yekta, A. J. Am. Chem. SOC.1978, 100, 5951. (15),Luo, H.; Boens, N.; van der Auweraer, M.; de Schryer, F. C.; Malliaris, A. J. Phys. Chem. 1989, 93, 3244.
Effect of Alcohols on Catalysis by Dodecyl Sulfate Micelles
The Journal of Physical Chemistry, Vol. 96, No. 4, 1992 1959
'I 'T
I
04
3
0
6
9
12
0 0
1
2
3
4
5
6
102[SDS]/M Figure 3. Influence of SDS concentration (=alcohol concentration) on the experimental rate constant ko for the acid denitrosation of MNTS in the presence of (0) 1-hexanol or (0) 1-butanol. [MNTSIo = 5.4 X M and [HCI] = 0.1 17 M. The solid lines were calculated by using the PIE model (see Table 111).
obtained for HDS were 55-58 with and without alcohol, even when the concentration of surfactant (=[alcohol]) was greater than that at which the second change of slope occurred in the conductivity-[HDS] plots. These values are typical of spherical micelles and rule out the possibility that the second change of slope was due to altered micellar shape. They are in keeping with Lianos et al.'sl' finding that concentrations of pentanol much greater than those used by us are required to cause significant changes in the shape and aggregation number of SDS micelles. An yanomalous" increase in specific conductivity at high surfactant concentrations has been reported previously in sporadic cases. Bunton et al.,I8 for example, found that the conductivity of alkanesulfonic acid micelles rose faster than surfactant concentration at concentrations greater than 0.01 M and suggested that this might be due to the micellar surface possibly conducting protons better than the aqueous phase. This hypothesis would explain why we observed no such effect with SDS. Be that as it may, since investigation of this phenomenon was not the primary aim of this work, we decided to calculate /3 by Hoffmann'si3 method using only conductivity values for HDS concentrations less than 4 X M. The resulting values of /3 (Table I) do not differ significantly from those found for the corresponding SDS-alcohol mixtures. Kinetics of tbe Denitrosation of MNTS in Wate~SDsAlcohol The influence of the concentrations of SDS and alcohol on the reaction rate was studied in the presence of a fixed concentration of HCl. Figures 3 and 4 show typical results, which are qualitatively the same as those obtained with SDS alone.IO The degree of catalysis in all cases increased to a maximum with surfactant/alcohol concentration and then decreased progressively. However, the maximum degree of catalysis attained was in general less than that achieved with SDS alone and decreased with increasing alcohol chain length. This characteristic pattern of behavior may be explained in terms of the PIE model.lI This model is based on Menger's pseudophase model,I9 which assumes that micelles are uniformly distributed throughout the solvent and treats them as a phase distinct from the solvent phase. It leads to the following expression ~
(16) Coll, H.J . Phys. Chem. 1970, 74, 520. (17) Lianos, P.; Lang, J.; Strazielle, C.; Zana, R. J . Phys. Chem. 1982,
86, 1019.
(18) Bunton, C. A.; Romsted, L. S.; Savelli, G. J . Am. Chem. SOC.1979, 101, 1253. (19) Menger, F. M.; Portnoy, C. A. J . Am. Chem. SOC.1967.89, 4698.
1 O 2 [SDS]/M
Figure 4. Influence of SDS concentration (=alcohol concentration) on the experimental rate constant ko for the acid denitrosation of MNTS in the presence of (0)benzyl alcohol or (0) tert-butyl alcohol. [MNTSIo = 5.4 X M and [HCI] = 0.117 M. The solid lines were calculated by using the PIE model (see Table 111).
for the pseudo-first-order rate constant of the reaction studied here:
where kwand kZmare the bimolecular reaction rate constants in the aqueous and micellar phases, respectively, P is the molar reaction volume per mole of micellized surfactant for reaction in the micellar phase, K, is the micellesubstrate binding constant, [D,] is the concentration of micellized surfactant ([D,] = [D] - cmc, where [D] is the total concentration of surfactant), and mH = [H+,]/[D,] is the concentration of H+ bound to the micelle, expressed as a mole ratio. If the reaction is assumed to take place in the Stem layer, the volume of the latter, uM,which is estimated as 0.14 M-1,18*20 can be taken as the molar reaction volume. The PIE model incorporates into the above theory the assumption that the surfactant's own counterion (in this case inert Nae) competes with the reactive cations (in this case H+) for binding positions in the Stem layer, giving rise to the ion exchange equilibrium
+
Na+w H+M 2 Na+M + H+w KNaH (4) where subscripts M and W refer to the micellar and aqueous phases, respectively. This equilibrium condition leads to the equation
dW+ltotal = 0 (5) (KNaH - iDnl For simple surfactant solutions, P i s constant in eq 3 and eqs 3 and 5 can be jointly fitted to experimental data to yield K,, kZm, and KNaH.The binding of alcohol, however, swells the micelles,b8 and to deal with this circumstance, it has been suggestedb8,*' that V be taken as the molar reaction volume of the unmixed micelle plus the molar volume of the bound alcohol: = P M + (%OH/nM)PRkoH (6) where (nROH/nM) is the mole ratio between the bound alcohol and the micellized surfactant and PRoH is the molar volume of pure ~
(20) (a) Bunton, C. A. Caral. Rev.-Sci. Eng. 1979, 20, 1 . (b) AI-Lohedan, H. A.; Bunton, C. A,; Romsted, L. S. J. Phys. Chem. 1981,85, 2123. (21) Martin, C. A.; McCann, P. M.; Ward, M. D.; Angelos, G. H.; Jaeger, D. A. J . Org. Chem. 1984, 49, 4392.
1960 The Journal of Physical Chemistry, Vol. 96, No. 4, 1992 TABLE 111 Rate and Equilibrium Constants Obtained by Fitting Eqs 3 and 5 Simultaneously to Experimental Dah for the Acid DenitrosPtion of MNTS in Water-SDS-Alcohol Mixtures with Equal Concentrations of SDS and Alcohol 103k2m, KS? KROH! M-l M-I s-1 h a H alcohol M-' ~~
c 1-butanol
isobutyl alcohol 2-butanol terr-butyl alcohol 1-pentanol terf-amyl alcohol 1- hexanol benzyl alcohol 22.
0.0915 0.0926 0.0926 0.0963 0.109 0.109 0.126 0.103
3.25 4.25
3.00 1.23 13.75 6.50 47.50 8.25
108 146 111 90 106 158 111 114 135
3.45 4.41 4.95 4.45 4.14 4.90 3.78 4.60 4.00
0.75 0.70 0.80 0.40 0.70 0.60 0.70 0.70 0.80
Bravo et al.
I
c9 -
r)
0
Calculated from molecular weight and density data. bReference 'Reference 10.
alcohol, which can be calculated from molecular weight and density data. Since for low additive fractionsZZthe alcohol is distributed between the aqueous and micellar pseudophases in accordance with the equation KROH = [ROHMl/ [ROHWl [Dn] (7) (the distribution constant changes sharply when the fraction of alcohol in the micellar phase is about eq 6 becomes (for equal concentrations of surfactant and alcohol) p = Y M + KROHIDlvROH/(l + KROH[DnI) (8) Values of KRoH for a large number of alcohols and SDS micelles have been measured by StilbsZ2and are listed in Table 111. The widely differing values of KRoH listed show that the reaction volume depends heavily on which alcohol is added (increasing the SDS/alcohol concentration to 0.1 M increases Yfrom 0.14 to just 0.16 M-' if 2-butanol is used but to 0.24 M-'if 1-hexanol is employed). We have kept Stilbs' criterion22also when defining K,, which is therefore defined in terms of D,, independently of the amount of solubilized alcohol. In fitting eqs 3 and 5 to the experimental - [D] data obtained for each alcohol, Y was not held constant but was calculated from eq 8 using the values of VRoH and KRoH listed in Table 111. For 8, the values listed in Table I were used, and for the cmc, the value of 1 X M was used, which was the lowest concentration of surfactant at which kinetic alteration was observed. The cmc value is lower than the one observed in the absence of HCl, due to the lowering effect of the electrolytes upon its value.z4 For k, we used the rate constant measured in water alone, 3.4 X M-l s-I; this choice of a fixed kw was justified by the finding that, in the absence of surfactant and the presence of concentrations of tert-amyl alcohol of up to 0.073 M (approximately the maximum alcohol concentration used in the main work), the alcohol reduced k, by no more than 10%. Sets of ko - [D] - Ydata were simultaneously fitted to eqs 3 and 5, and values of K,, k2"',and KNaH were obtained for the different alcohols. Figures 3 and 4 illustrate the good fit obtained for all the alcohols used between the experimental data and the fitted equations. Table I11 lists the optimized values of K,, k2"',and KNaHtogether with those obtained in the absence of alcohol.'0 All three components are virtually independent of the identity of the alcohol; K, and k2"' are in general just slightly greater than the values obtained in the absence of alcohol, and the values of KNaH also agree fairly well with values obtained by direct measurement in the absence of alcohol.25 Thus none of the kinetic parameters depend to any great extent on the presence of alcohol, which must therefore affect the experimental rate constant solely through its (22) Stilbs, P. J. Colloid Interface Sci. 1982, 87, 385. (23) Abuin, E. A.; Lissi, E. A. J. Colloid Interface Sci. 1983, 95, 198. (24) Ptrez-Benito, E.; Rodenas, E. J. Colloid Inferface Sci. 1990, 139, 87. (25) (a) Quina, F. H.; Politi, M. J.; Cuccovia, I. M.;Martins-Franchetti, S. M.; Chaimovich, H. In Solution Behaviour of Surfactants: Theoretical and Applied Aspects; Mittal, K. L., Fendler, E. J., Eds.; Plenum Press: New York, 1982; Vol2, p 1125. (b) Bunton. C. A,; Ohmenzetter, K.; Sepulveda, L. J. Phys. Chem. 1977,81,2000. (c) Hafiane, A.; Issid, I.; Lemordant, D. J. Colloid Interface Sci. 1991, 142, 167.
r
0
2
4
6
1 O2 [HDS]/M
Figure 5. Influence of HDS concentration (=alcohol concentration)on ko for the acid denitrosation of MNTS in the presence of ( 0 )benzyl alcohol and (0)fert-butyl alcohol. [HCI]= 0.117 M. The solid lines
were calculated from the PIE model (see Table IV). increasing the reaction volume, i.e. by diluting the reagents in the micellar pseudophase. The greater the hydrophobicity and molar volume of the alcohol, the greater are the increase in reaction volume and the reduction in reaction rate. The good fit of the experimental data to the kinetic model and the finding that rate constants,binding constants, and ion-exchange constants are virtually independent of the nature of the added alcohol suggest that addition of alcohols to the micelles produce no significant change in the reaction medium. This is in keeping with the hypothesis5 that addition of alcohol is accompanied by water penetration into the micelles, so that the "real" polarity of the reaction medium would hardly be affected. In particular, the constancy of KNaHsuggests that the solvation of Na+ and H+,the latter especially, is unaffected by the presence of alcohol, therefore indicating a similar hydration shell with and without alcohol; i.e. the aqueous content of the Stern layer is not significantly decreased. This appears to contradict reports4 that the neighborhood of fluorescent probes is less polar in the presence of alcohol than in its absence. Since other studies26have found that the polarity of a fluorescent probe in the presence of alcohol is independent of the identity or hydrophobicity of the alcohol, it appears that the interpretation of fluorescent probe experiments may have to take into account both the composition and structural organization of the ~ o l v e n t . ~ Kioetics of the Deaitrosation of MNTS io Water-HDSAlcoM. Figure 5 shows typical results obtained with HDSalcohol micelles under the same experimental conditions as for the SDS study. HDS is stable under our working conditions, and its hydrolysis does not interfere with any of our meas~rements.~'As with SDS, the degree of catalysis depends on which alcohol was added. Unlike the Figures 3 and 4 plots for SDSalcohol micelles, those of Figure 5 show that, with HDS, catalysis tended to a maximum limiting value with increasing surfactant/alcohol concentration, but this too is in agreement with the PIE model, since in the absence of a competing ion mH = 8 (a constant) in eq 3, which for high surfactant concentrations then reduces to
ko = kzm@/Y (9) In fitting eq 3 (with mH = 8) to the experimental data, the values of KRoH required for calculation of P (eq 8) were, in the absence of published values for HDS-alcohol mixtures, assumed to be close to those used previously for binding to SDS micelles (Table 111), which were accordingly used here. This assumption (26) Lianos, P.; Lang, J.; Zana, R.J. Phys. Chem. 1982,86,4809. (27) (a) Kurz, J. L. J. Phys. Chem. 1962, 66, 2239. (b) Garnett, C. J.; Lambie, A. J.: Beck,W.H.; Liler, M . J. Chem. Soc., Faraday Trans. I 1983, 79, 953.
J. Phys. Chem. 1992, 96, 1961-1967 TABLE Iv: Rate and Equilibrium Coostants Obtained by Fitting Eq 3 (with m H = @)to Experimental Data for the Acid Denitroslltion of MNTS in Water-HDS-Alcohol Mixtures with Equal Concentrations of HDS and Alcohol
alcohol a
1-butanol isobutyl alcohol 2-butanol tert-butyl alcohol 1-pentanol tert-amyl alcohol 1-hexanol benzyl alcohol
K.. M-l 97 123 117
100 85 125 74 115 107
103w, M-l 3.78 3.66 4.69 4.41 4.34 3.64 4.55 5.62 3.82
s-1
Reference 10. was based on the fact that HDS and SDS have the same hydrocarbon chain, an important influence on the binding of amphiphiles to micelles,3ethe similarity between the measured aggregation numbers (and hence shape) of SDS-alcohol and HDSalcohol micelles (Table 11), and the similarity between the behaviors of HDS and SDS micelles in catalyzing the acid denitrosation of MNTS in the absence of alcohol (in particular, they both have the same affinity for MNTS and the same KNaHvalue).Io Table IV lists the values of K, and k2” obtained using an optimization program based on Marquardt’s2* algorithm, with
1961
values of /3 experimentally measured (Table I) and values of k, and the cmc the same as those for SDS. Both constants vary very little with the alcohol used and are very similar to the values obtained with SDS, so that again the effect of the alcohol is merely to increase the volume of reaction. This similarity between the two kinds of mixed micelles is quite as expected, since the two surfactants differ only as regards their counterion and also behave almost identically in the absence of alcohol.I0 The assumption that KRoHhas very similar values for SDS and HDS is supported by these quite reasonable findings. In conclusion, the addition of moderate quantities of mediumchain alcohols to anionic SDS and HDS micelles caused a reduction in reaction rate that can be explained as due purely to the resulting dilution of the reagents, with no change in kinetic or thermodynamic constants occurring.
Acknowledgment. C.B. thanks the Xunta de Galicia for the research training grant allowing him to take part in this research. Financial support from the Xunta de Galicia (Project XUGA 20903A90) is gratefully acknowledged. We also thank the referees for their comments. Rdstry NO. MNTS,80-11-5; HDS, 1510-16-3;SDS, 151-21-3; l-butanol, 71-36-3; 2-butanol, 78-92-2; isobutyl alcohol, 78-83-1; fertbutyl alcohol, 75-65-0; 1-pentanol,71-41-0; tert-amyl alcohol, 75-85-4; 1-hexanol, 11 1-27-3;benzyl alcohol, 100-51-6. (28) Marquardt, D. W. J . SOC.Ind. Appl. Math. 1963, 1 1 , 431.
Lanthanide Shift NMR Studies of Bile Salt Aggregates Steven M. Meyerhoffer: Thomas J. Wenze1,t and Linda B. McGown*Yt Department of Chemistry, P. M. Gross Chemical Laboratory. Duke University, Durham, North Carolina 27706, and Department of Chemistry, Bates College, Lewiston, Maine 04240 (Received: August 26, 1991; In Final Form: October IO, 1991)
The binding of Tb3+,Dy”, Eu”, Tm3+, and Yb3+ with sodium taurocholate (NaTC), 3-[(3-~holamidopropyl)dimethylammonio]-2-hydroxy-1-propanesulfonate (CHAPSO), and 3-[(3-~holamidopropyl)dimethylammonio] - 1-propanesulfonate (CHAPS) was studied by ’H NMR spectroscopy. Proton assignments are given for CHAPSO and CHAPS, and the proper selection of an internal reference for NMR shift studies is addressed. Results for NaTC indicate that the lanthanide ion binds asymmetrically with the anionic sulfonate group and the oxygen atom of the peptide carbonyl through a bidentate interaction. Association constants and bound shifts were calculated from the ‘H NMR shift data using a nonlinear least-squares regression analysis. In all cases, the shift data were fit to an equilibrium assuming 1:l lanthanidetaurocholate complexation. In micellar NaTC, however, the association constants calculated with the onestep model only approximate the overall binding interaction; the equilibrium may be influenced by the formation of a 1:4 stoichiometric complex as well as by the total concentrationof lanthanide salt present in solution. The magnitudes of the observed shifts and calculated association constants for Tb3+-taurocholate complexes were greater in micellar NaTC than in premicellar NaTC, indicating an increased binding strength in micellar solution. The Tb3+ion was also observed to bind with the anionic sulfonate group of CHAPSO and CHAPS, but the interactions were significantly weaker than with the taurocholate anion.
Introduction The bile salts are naturally occurring detergents which form micellar aggregates in aqueous solution.’ In addition to their important physiological role? bile salts have also been studied in recent years as alternatives to conventional detergents for chemical analysis. Specific areas of application include chemical separat i o n ~and ~ luminescence analy~is.~Bile salt micelles are smaller and more rigid than those of conventional detergents, resulting in unique aggregation behavior with respect to self-association as well as solubilization of hydrophobic molecules in aqueous soluti~n.~-~ *Author to whom correspondence should be addressed. ‘Duke University. *Bates College.
The effects of metal ions on the fluorescence intensity of molecular probes in aqueous solutions of the trihydroxy bile salt sodium taurocholate (NaTC)8,9and of 3-[(3-cholamidopropyl)dimethylammonio] -2-hydroxy- 1-propanesulfonate (CHAPSO)? (1) Small, D. M. In Bile Acids; Nair, P. P., Kritchevsky, D., Eds.; Plenum Press: New York, 1971; pp 249-357. (2) Hofmann, A. F.; Small, D. M. Annu. Reu. Med. 1967, 18, 333. (3) Buckley, J. J.; Wetlaufer, D. B. J . Chromnrogr. 1989, 464, 61. (4) McGown, L. B.; Kreiss, D. S. Fluorescence Detection 11. Proc. SPIE-Int. SOC.Opt. Eng. 1988, 910, 73-80. (5) Zana, R.; Guveli, G. J . Phys. Chem. 1985,89, 1687. (6) Mukerjee, P.; Moroi, Y.;Murata, M.; Yang, A. Y. S. Hepatology 1984,4, 61s. (7) Fisher, L.; Oakenfull, D. Aust. J . Chem. 1979, 32, 31. (8) Nithipatikom, K.; McGown, L. B. A n d . Chem. 1988, 60, 1043. (9) Meyerhoffer, S. M.; McGown, L. B. J . Am. Chem. SOC.1991, 113, 2 146.
0022-3654/92/2096-1961$03.00/00 1992 American Chemical Society
