J. Phys. Chem. B 2007, 111, 12185-12194
12185
Kinetic Coupled with UV Spectral Evidence for Near-Irreversible Nonionic Micellar Binding of N-Benzylphthalimide under the Typical Reaction Conditions: An Observation Against a Major Assumption of the Pseudophase Micellar Model May Ye Cheong, Azhar Ariffin, and M. Niyaz Khan* Department of Chemistry, Faculty of Science, UniVersity of Malaya, 50603 Kuala Lumpur, Malaysia ReceiVed: May 23, 2007; In Final Form: August 7, 2007
Pseudo-first-order rate constants (kobs) for alkaline hydrolysis of N-benzylphthalimide (1) show a nonlinear decrease with the increase in [CmEn]T (total concentration of Brij 58, m ) 16, n ) 20 and Brij 56, m ) 16, n ) 10) at constant [CH3CN] and [NaOH]. These nonionic micellar effects, within the certain typical reaction conditions, have been explained in terms of the pseudophase micellar (PM) model. The values of micellar binding constants (KS) of 1 are 1.04 × 103 M-1 (at 1.0 × 10-3 M NaOH) and 1.08 × 103 M-1 (at 2.0 × 10-3 M NaOH) for C16E20 as well as 600 M-1 (at 7.6 × 10-4 M NaOH) and 670 M-1 (at 1.0 × 10-3 M NaOH) for C16E10 micelles. The pseudo-first-order rate constants (kM) for hydrolysis of 1 in C16E20 micellar pseudophase are ∼90-fold smaller than those (kW) in water phase. The values of kM for hydrolysis of 1 in C16E10 micelles are almost zero. Kinetic coupled with UV spectral data reveals significant irreversible nonionic micellar binding of 1 molecules in the micellar environment of nearly zero hydroxide ion concentration at g0.14 M C16E20 and 1.0 × 10-3 M NaOH while such observations could not be detected at e0.17 M C16E20 and 2.0 × 10-3 M NaOH. Significantly, such irreversible C16E10 micellar binding of 1 molecules could be detected at 8.8 × 10-2 M C16E10 and 1.0 × 10-3 M NaOH as well as at g3 × 10-3 M C16E10 and 7.6 × 10-4 M NaOH, while the rate of hydrolysis of 1 is completely ceased at g0.05 M C16E10 and 7.6 × 10-4 M NaOH. The rate of hydrolysis of 1 at 5.0 × 10-2 and 8.8 × 10-2 M C16E10 and 1.0 × 10-3 M NaOH reveals the formation of presumably phthalic anhydride, whereas such observation was not observed in the C16E20 micellar system under similar experimental conditions.
Introduction
SCHEME 1
The classical pseudophase micellar (PM) model and its various extended forms contain a basic and common assumption that the micellar reaction environment/solubilization site constitutes a uniform and homogeneous medium.1 But indirect experimental evidence is making the fact increasingly clearer that the micellar reaction/solubilization environment is not completely homogeneous in terms of polarity/dielectric constant, water concentration, water structure, viscosity, and ionic strength (for ionic micelles only).2 But, it seems that the changes in these physicochemical properties with change in distance from the exterior to deep interior of the micelle are gradual and continuous.3 Perhaps because of this micellar characteristic, the multiple micellar pseudophase (MMP) model of Davies led to the kinetic equation similar (in form) to one derived based upon the classical two-state PM model with a modified definition of micellar binding constant of reactants and rate constant for micellar-mediated reaction.4 There appears to be no report in the literature against this generalization. However, irreversible micellar portioning of two different reactants in terms of hydrophilicity has been recently observed at larger than a typical value of [C12E23]T/[NaOH] () Rt) in C12E23 micellar-mediated alkaline hydrolysis of phthalimide,5 phenyl benzoate (PB),6 and phenyl salicylate (PSH)6 as shown in reaction Scheme 1. These observations cannot be ascribed to a shielding effect of micelles because the observed data, obtained at R () [C12E23]T/[NaOH]) values below the typical Rt value, have been explained in terms * Corresponding author. E-mail:
[email protected].
of the PM model while shielding the effect of the micelles is not expected to display such micellar-mediated reaction characteristics. The shielding effect of micelles depends only upon the difference of the hydrophobicity of the two different reactant molecules, and it should be independent of the nature of micelleforming surfactants. Irreversible micellar trapping of PSH7 and PB8 could not be detected in pH-independent hydrolysis and hydrazinolysis of PSH in the presence of anionic micelles and alkaline hydrolysis of PB in the presence of C16E20, micelles. These observations are not in favor of a shielding effect as the cause of the irreversible trapping of PSH and PB at R values >Rt.5
10.1021/jp073980j CCC: $37.00 © 2007 American Chemical Society Published on Web 10/03/2007
12186 J. Phys. Chem. B, Vol. 111, No. 42, 2007
Cheong et al.
But such unusual observations were not obtained in the presence of C16E20 micelles under essentially similar conditions.5,8 The observed data, obtained at R < Rt, fit satisfactorily to the PM model of the micelle. Perhaps the most unusual observation, revealed by recent studies,5,6,9 is an apparent depletion of hydroxide ions and water molecules from C12E23 micellar environment where organic substrate molecules (sub ) PSH, PB, phthalimide, and 4-nitrophthalimide) reside, and consequently, it causes irreversible trapping of sub molecules by micelles at R values >Rt. Such irreversible micellar trapping of sub molecules is expected to increase with the increase in hydrophobicity of sub molecules. The present study was initiated to explore the possibility of micellar irreversible portioning of reactants HO- and N-benzylphthalimide (1), a nonionizable and more hydrophobic imide compared with phthalimide, in the C16E20 and C16E10 micellar-mediated alkaline hydrolysis of 1.
The observed data and their plausible explanations are described in this manuscript. Experimental Section Materials. Reagent-grade polyoxyethylene (20) cetyl ether {C16H33(OCH2CH2)20OH (Brij 58 or C16E20)} and polyoxyethylene (10) cetyl ether {C16H33(OCH2CH2)10OH (Brij 56 or C16E10)} were commercial products of highest available purity. All other chemicals were also of reagent-grade. NBenzylphthalimide (1) and N-benzylphthalamic acid (2) were synthesized as described in an earlier reports.10 Stock solutions (0.01 M) of 1 were prepared in acetonitrile. Kinetic Measurements. The rate of hydrolysis of 1, in an alkaline medium, was studied by monitoring the disappearance of reactant 1 spectrophotometrically at 300 nm. In a typical kinetic run, the reaction mixture (total volume of either 4.9 cm3 for fast reactions with half-life of 5 h) containing all the reaction ingredients except 1 was equilibrated at 35 °C (using a thermostated water bath) for 5-10 min. The reaction was then started by adding 0.1 cm3 (to 4.9 cm3) or 0.3 cm3 (to 14.7 cm3) of reaction mixture of 0.01 M 1 (prepared in CH3CN). The final total volume of the reaction mixture in each kinetic run was 5 cm3 (for fast reaction) or 15 cm3 (for slow reaction). An aliquot of ∼2.5 cm3 was quickly transferred to a 3 cm3 quartz cuvette kept in the thermostated cell compartment of the UV-vis double-beam spectrophotometer. The decrease in absorbance (Aobs) at 300 nm as a function of reaction time (t) for fast reaction was monitored by the spectrophotometer. A sampling technique was used to monitor the disappearance of 1 spectrophotometrically at 300 nm as a function of reaction time (t) for slow reactions. All the kinetic runs were carried out under essentially pseudofirst-order kinetic conditions, and pseudo-first-order rate constants (kobs) were calculated from eq 1
Aobs ) δapp[R0] exp(-kobs t) + A∞
(1)
with nonlinear least-squares technique considering δapp (apparent molar absorptivity of the reaction mixture) and A∞ (the absorbance of the reaction mixture at reaction time t ) ∞) also as unknown kinetic parameters.11 In eq 1, Aobs represents absor-
bance of the reaction mixture at any reaction time t and [R0] is the initial concentration of 1. The reactions were carried out for up to 7-9 half-lives, and the observed data (Aobs vs t) fitted well to eq 1 as is evident from low values of standard deviations associated with the calculated kinetic parameters, kobs, δapp, and A∞ (Tables S1-S4, Supporting Information) as well as from the least-squares predicted solid lines of Figure 5. The molar absorptivities of ionized and nonionized N-benzylphthalamic and phthalic acids as well as benzylamine at 300 nm are very low (∼40 M-1 cm-1). Therefore, δapp ≈ δ1 at 300 nm because δapp ) δ1 - δP , where δ1 and δP represent the molar absorptivity of 1 and product (N-benzylphthalamic acid, 2), respectively. Most of the kinetic runs were carried out at low hydroxide ion concentration (7.6 × 10-4 to 2.0 × 10-3 M NaOH) simply because of the rather high value of the second-order rate constant (kOH ) 22 M-1 s-1 at 35 °C) for hydroxide ion-assisted hydrolysis of 1.10 However, we did not use any special device to protect reaction mixtures from atmospheric CO2 except usual ones: using g5 M NaOH stock solution and freshly glassdistilled water. The stock solutions of lower hydroxide ion concentrations were prepared in freshly glass-distilled water. Reactions were carried out in properly stoppered reaction vessels and cuvettes. The average values of kOH (obtained from g5 kinetic runs) at 7.6 × 10-4, 1.0 × 10-3, and 2.0 × 10-3 M NaOH as well as within the [CmEn]T (m ) 16, n ) 20, 10) range from 0.0 to 0.98 at 0.088 M C16E10. Thus, these results show that ∼34% micelle-bound 1 hydrolyzed through the usual equilibrium process 1W + Dn h 1M assumed in the PM model while ∼64% micelle-bound 1
Figure 5. Plots of Aobs vs reaction time for hydrolysis of NBPT at 5.0 × 10-2 M {([) for first set and (2) for second set} and 8.8 × 10-2 M {(9) for first set and (b) for second set} [C16E10]T at 1.0 × 10-3 M NaOH. Solid lines are drawn through the least-squares calculated data points.
remained nearly irreversibly trapped by micelles. These results can best be represented by Scheme 5 where n1 1M and n2 1M are in equilibrium with n 1W while n3 1M and n4 1M are not in equilibrium with n 1W. Similarly, s1 HO-M and s2 HO-M ions are in equilibrium with s HO-W ions. Unexpected Minima in the Plots of Aobs versus t for C16E10Mediated Alkaline Hydrolysis of 1 Under the Typical Reaction Conditions. The kinetics of the hydrolysis of 1 at 5.0 × 10-2 and 8.8 × 10-2 M [C16E10]T in the presence of 1.0 × 10-3 M NaOH was complicated by an unexpected increase in the Aobs with the increase in reaction time, t, at >21 and 76 h for the reaction rate studied at 5.0 × 10-2 and 8.8 × 10-2 M [C16E10]T, respectively. The representative plots of Aobs versus reaction time (t) are shown in Figure 5. These plots show a rapid monotonic decrease in Aobs in the initial phase of the reaction followed by a slow monotonic increase with increasing reaction time in the latter phase of the reaction at 300 nm. Similar observations have been reported in the reaction of HOwith phthalimide in the presence of g1.2 × 10-2 M C12E23 in mixed aqueous solvent containing 2% v/v CH3CN, 2.0 × 10-2 M NaOH, 1.0 × 10-2 M CTABr, and 2.0 × 10-4 M phthalimide.5 The minima in the plots of Figure 5 demonstrate the formation of a stable intermediate product which absorbs significantly at 300 nm. The most probable such a stable intermediate is phthalic anhydride (3) because δ3 ≈ 2300 M-1 cm-1, δ2 ) δ5 ) 40 M-1 cm-1, and δ4 ) 4 M-1 cm-1 at 300 nm where 2, 4, and 5 represent N-benzylphthalamic acid, benzylamine, and phthalic acid, respectively. At pH e 3, the rate of N-cyclization (rate constant, kobs ) 1.25 × 10-6 s-1) is almost negligible compared to that of O-cyclization (rate constant, kobs ) 17.6 × 10-6 s-1) in the
Nonionic Micellar Binding of N-Benzylphthalimide
J. Phys. Chem. B, Vol. 111, No. 42, 2007 12193
TABLE 4: Values of Kinetic Parameters, k1, k2, and A∞ Calculated from Eq 6a 105 k1 s-1
107 k2 s-1
102 A∞
103∑di2 e
0 0
34.3 ( 1.7f 34.0 ( 1.7
1.88 ( 0.09f 2.51 ( 0.14
8.2 ( 0.4f 8.1 ( 0.4
2.193g 2.258
0 0
35.5 ( 1.3 35.2 ( 1.4
1.97 ( 0.08 2.49 ( 0.11
6.7 ( 0.3 6.6 ( 0.3
1.633 1.749
800 630
0.65 0.65
8.74 ( 1.70 8.76 ( 1.74
1.34 ( 0.26 1.78 ( 0.38
34.1 ( 0.4 34.1 ( 0.5
2.110g 2.169
800 730 550
0.65 0.65 0.65
11.9 ( 1.8 11.5 ( 1.7 10.6 ( 1.6
3.27 ( 0.43 3.91 ( 0.55 7.38 ( 1.42
30.5 ( 0.4 30.4 ( 0.4 30.2 ( 0.4
2.202h 2.105 1.867
[C16E10]T M
δ1app M-1 cm-1 b
δ3app M-1 cm-1 c
0.050
1745 1745
2300 1860
0.050
1750 1750
2300 1940
0.088
645 645
0.088
687 687 687
FIT1d
a [10] ) 2.0 × 10-4 M, 35 °C, λ ) 300 nm, [NaOH] ) 1.0 × 10-3 M. b The values of δ1app were obtained from Table S4 (Supporting Information), where δapp ) δ1app. c δ3app ) (1- FIT1)δ30 with δ30 ) 2300 M-1 cm-1. d The values of FIT1 were calculated from eq 4 with values of A∞, AT, AT+2, and A0 from Table S4 (Supporting Information). e di ) Aobs,i - Acalcd,i where the values of Acalcd,i at the ith reaction time, ti, were obtained from eq 6 with kinetic parameters listed in Table 4. f Error limits are standard deviations. g Reaction mixture contained 2% v/v CH3CN. h Reaction mixture contained 0.67% CH3CN.
cleavage of 2 in aqueous reaction mixture containing 2% v/v CH3CN.24 Thus, the aqueous cleavage of 1 follows an irreversible reaction path: k1
k2
k3
1 98 2 98 3 + 4 98 4 + 5
(5)
The respective absence and presence of a minimum at [C16E10]T e 3.0 × 10-2 M and at [C16E10]T g 5.0 × 10-2 M may be ascribed to the consequence of the effects of [C16E10]T on the pH of the micellar environment where 2 molecules reside (i.e., 2M molecules with subscript M representing micellar pseudophase). It is obvious that at [C16E10]T e 3.0 × 10-2 M, the pH of the micellar reaction environment of 2M remained considerably high and hence the conversion of 2 to 3 in the k2 step was completely stopped because, in view of related studies,14,25 anionic N-benzylphthalamic acid (2-) did not undergo O-cyclization to give 3. However, at [C16E10]T g 5.0 × 10-2 M, the pH of the micellar reaction environment of 2M dropped to a level where there was significant amount of nonionized 2M which caused the kinetically detectable occurrence of the conversion of 2 to 3. An increase in [C16E10]T at a constant [NaOH] perhaps decreases both [H2O] and [HO-] in the micellar region of 2M molecules, which in turn increases the fraction of nonionized 2M. The increase in the fraction of nonionized 2M would increase the rate of formation of 3. Although the data are not sufficient to carry out a detailed kinetic analysis of the formation of 3, an attempt was made to fit the observed data at [C16E10]T g 5.0 × 10-2 M (as shown in Figure 5, where there is a maximum change in absorbance in the final phase of the reaction) to eq 6 which is derived from the reaction scheme in eq 5 with the conceivable conditions that k3 , k2 and δ2 , δ1 and δ3 at 300 nm. The assumption that k3 , k2 is based upon the mere observation of increase in Aobs with increasing t until the last observed Aobs value at t ) 983 h. This shows that the water concentration in the micellar environment of reaction 2M f 3M + 4M, i.e., the k2 step in eq 5, is extremely low. The plots in Figure 5 should have resulted in maxima if k2 ≈ k3 or k2/k3 ≈ 0.5-2.0 because δ4 + δ5 (≈45 M-1 cm-1 at 300 nm) , δ3 () 2300 M-1 cm-1 at 300 nm). Similarly the plots in Figure 5 should not reveal minima or maxima provided k3 . k2. In eq 6, [R0] represents the initial concentration of 1 () 2.0 × 10-4 M), δ1app ) δ1 - δ2, δ3app ) δ3 - δ2, and A∞ ) δ2[R0] provided δ4 ≈ 0.
Aobs ) δ1app[R0] exp(-k1t) + δ3app[R0] × 1 1+ [k exp(-k1t) - k1 exp(-k2t)] + A∞ (6) k 1 - k2 2
{
}
The nonlinear least-squares fit of the observed data (Figure 5) to eq 6 was carried out by considering k1, k2, and A∞ as unknown parameters at the given values of δ1app and δ3app. The mathematical derivation of eq 6 is described in the Supporting Information. The decrease in δ1app with increase in [C16E10]T in the initial phase of the plots in Figure 5 is due to nonreacted 1 molecules trapped irreversibly in the micellar region where [HO-M] ≈ 0. The value of δ1app was obtained from eq 1 (where δapp ) δ1app) by using the observed data of the fast initial phase of the reaction. The value of δ3app was obtained from the relationship δ3app ) (1 - FIT1)δ30 where FIT1 represents the mole fraction of 1 that remained trapped by micelles in a micellar environment of nearly zero hydroxide ion concentration and δ30 is the molar extinction coefficient of 3 at 300 nm in an aqueous solvent. The values of FIT1 at the typical values of [C16E10]T were calculated from eq 4. The respective values of δ1 and δ3 at 300 nm vary from 2250 to 1680 M-1 cm-1 and 2300-2000 M-1 cm-1 with increase in acetonitrile content from 2% to 90% in mixed aqueous solvent. However, the values of δ2 of ∼30 M-1 cm-1 at 300 nm remained unchanged with change in acetonitrile content from 2% to 90% v/v in mixed aqueous solvent containing 0.02 M HCl. An apparent satisfactory fit of observed data to eq 6 is evident from the plots of Figure 5, where solid lines are drawn through the least-squares calculated data points. The nonlinear least-squares calculated values of k1, k2, and A∞ are shown in Table 4. The values of k1, k2, and A∞ were calculated at different presumed values of δ3app, because of the uncertainty in the water concentration of micellar reaction environment. However, the results summarized in Table 4 reveal that a significant change in the value of δ3app caused almost no change in k1 values and a slight change in ∑di2, whereas the value of k2 was significantly affected. The calculated values of k1 are similar to the corresponding values of kobs (Table S4, Supporting Information) calculated from eq 1 using Aobs values within the reaction time (t) range of 0.02-21 h at 0.050 M C16E10 and 0.03-31 as well as 76 h at 0.088 M C16E10. The values of k2 at 0.050 and 0.088 M C16E10 are nearly 102-fold smaller than the corresponding value of 2.44 × 10-5 s-1 obtained in the absence of micelles at 0.02 M HCl and 90% v/v CH3CN in mixed aqueous solvent.24 These results may be largely ascribed to the presence of the extremely low value of the fraction of nonionized 2M, F{) [2M]/([2M] + [2-M]), where 2-M represents ionized 2M}.
12194 J. Phys. Chem. B, Vol. 111, No. 42, 2007 Conclusions It appears that the observed data (Aobs vs t) for the rate of alkaline hydrolysis of 1 in the presence of C16E20 micelles follow the PM model until a typical value of Rt ) R ) [C16E20]T/ [NaOH] where the values of [C16E20]T were varied at a constant value of [NaOH]. The observed data could not follow the PM model at R > Rt. Similar observations were obtained in the presence of C16E10 micelles. The value of Rt seems to depend upon both the nature of nonionic surfactant and the value of [CmEn]T/[NaOH] () R). The most unusual observation is the sharp decrease in water and hydroxide ion concentrations in the micellar environment of 1M molecules and consequently irreversible micellar trapping of 1 molecules at R > Rt. The alkaline hydrolysis product 2 of 1 is apparently more polar than 1, and at certain typical values of R for C16E10 micelles, the pH of the micellar environment of 2M became < ∼4 because, under such conditions, 2M molecules underwent O-cyclization producing 3 as the final stable product. Such observations were not detected in the presence of C16E20 micelles although irreversible C16E20 micellar trapping of 1 molecules has been detected at certain values of R. Differences in the behaviors of the two surfactants could be most likely related to the differences in the structures. However, the reason related to reactive or higher relative impurity in one of the two surfactants (C16E20 and C16E10) cannot be completely ruled out. Acknowledgment. The authors thank the National Scientific Research and Development Council of Malaysia for IRPA (Grant No. 09-02-03-0147), the Academy Sciences Malaysia for SAGA (Grant No. 66-02-03-0039), and the University Malaya for financial assistance. Supporting Information Available: Tables S1-S4 and mathematical derivation of eq 6 of the manuscript. This material is available free of charge via the Internet at http://pubs.acs.org. References and Notes (1) (a) Bunton, C. A. Catal. ReV.sSci. Eng. 1979, 20, 1. (b) Bunton, C. A.; Savelli, G. AdV. Phys. Org. Chem. 1986, 22, 213. (c) Bunton, C. A.; Nome, F.; Quina, F. H.; Romsted, L. S. Acc. Chem. Res. 1991, 24, 357.
Cheong et al. (2) Khan, M. N. Micellar Catalysis. In Surfactant Science Series; CRC Press, Taylor & Francis Group: Boca Raton, FL, 2006; vol. 133, p 33 and references therein. (3) (a) Mishra, A.; Behera, R. K.; Behera, P. K.; Mishra, B. K.; Behera, G. B. Chem. ReV. 2000, 100, 1973. (b) Mishra, A.; Patel, S.; Behera, R. K.; Mishra, B. K.; Behera, G. B. Bull. Chem. Soc. Jpn. 1995, 70, 2913. (c) Laschewsky, A. Curr. Opin. Colloid Interface Sci. 2003, 8, 274. (d) Raghuraman, S.; Pradhan, S. K.; Chattopadhyay, A. J. Phys. Chem. B 2004, 108, 2489. (e) Sterpone, F.; Pierleoni, C.; Briganti, G.; Marchi, M. Langmuir 2004, 20, 4311. (4) Davies, D. M.; Gillitt, N. D.; Paradis, P. M. J. Chem. Soc., Perkin Trans. 2 1996, 659. (5) Khan, M. N.; Ismail, E. J. Phys. Org. Chem. 2002, 15, 374. (6) Khan, M. N.; Ismail, E.; Yusoff, M. R. J. Phys. Org. Chem. 2001, 14, 669. (7) Khan, M. N. J. Chem. Soc., Perkin Trans. 2 1990, 445. (8) Khan, M. N.; Ismail, E. J. Phys. Org. Chem. 2004, 17, 376. (9) (a) Khan, M. N.; Ariffin, Z.; Yusoff, M. R.; Ismail, E. J. Colloid Interface Sci. 1999, 220, 474. (b) Khan, M. N.; Ismail, E. J. Colloid Interface Sci. 2001, 240, 636. (10) Cheong, M. Y.; Ariffin, A.; Khan, M. N. Indian J. Chem. 2005, 44A, 2055. (11) Khan, M. N. J. Chem. Soc., Perkin Trans. 2 1990, 435. (12) Broxton, T. J.; Duddy, N. W. Aust. J. Chem. 1979, 32, 1717. (13) Bunton, C. A.; Nayak, B.; O’Connor, C. J. Org. Chem. 1968, 33, 572. (14) Ariffin, A.; Khan, M. N. Bull. Korean Chem. Soc. 2005, 26, 1037. (15) Khan, M. N. Int. J. Chem. Kinet. 1987, 19, 143. (16) (a) Menger, F. M.; Portnoy, C. E. J. Am. Chem. Soc. 1967, 89, 4698. (b) Martinek, K.; Yatsimirskii, A. K.; Levashov, A. V.; Berezin, I. V. In Micellization, Solubilization, Microemulsion; Mittal, K. L., Ed.; Plenum: New York, 1977; Vol. 2, p 489. (17) (a) Bunton, C. A. Encyclopedia of Surface and Colloid Science; Hubbert, A. T., Ed.; 2002; p 980. (b) Bunton, C. A. J. Mol. Liq. 1997, 72, 231. (c) Cordes, E. H.; Gitler, C. Prog. Bioorg. Chem. 1973, 2, 1. (d) Bunton, C. A. In Surfactant in Solutions; Mittal, K. L., Shah, D. O., Eds.; Plenum: New York, 1991; Vol. 11, p 17. (18) Khan, M. N. Micellar Catalysis. In Surfactant Science Series; CRC Press, Taylor & Francis Group: Boca Raton, FL, 2006; vol. 133, pp 215216. (19) Broxton, T. J.; Wright, S. J. Org. Chem. 1986, 51, 2965. (20) Khan, M. N.; Ariffin, Z. Langmuir 1996, 12, 261. (21) Huibers, P. D. T.; Lobanov, V. S.; Katrizky, A. R.; Shah, D. O.; Karelson, M. Langmuir 1996, 12, 1462. (22) Rispens, T.; Engberts, J. B. F. N. J. Org. Chem. 2003, 68, 8520. (23) Lajis, N. H.; Khan, M. N. J. Phys. Org. Chem. 1998, 11, 209. (24) Cheong, M. Y.; Ariffin, A.; Khan, M. N. Submitted for publication. (25) (a) Hawkins, M. D. J. Chem. Soc., Perkin Trans. 2 1976, 642. (b) Blackburn, R. A. M.; Capon, B.; McRitchie, A. C. Bioorg. Chem. 1977, 6, 71.
