J. Chem. SOC., Faraday Trans. I, 1989, 85(3), 521-535 Acid-Base Equilibria in Aqueous Micellar Solutions Part 1 .-' Simple' Weak Acids and Bases Calum J. Drummond,*T Franz Grieser and Thomas W. Healy Department of Physical Chemistry, The University of Melbourne, Parkville, Victoria, 3052, Australia The acid-base equilibria of a number of phenols, amines and carboxylic acids in aqueous micellar solutions and organic solvent-water mixtures have been examined. For the majority of the molecules investigated, the differences between the plr$ values in pure water and the apparent plr$ values when the molecules reside within micellar interfacial microenvironments can primarily be ascribed to the differences between the mean intrinsic solvent properties of the interfacial and bulk phases, with an additional contribution from the electrostatic micellar surface potential in the case of the charged aqueous micellar solutions.The apparent p& of a weak acid or weak base residing at or in the vicinity of a charged interface, pKtbs, is generally considered to be composed of an electrostatic component due to the surface potential and an inherent interfacial non-electrostatic component. This relationship is often expressed as1-26 where pK," is the apparent p& of the molecule in the absence of any potential, Y is the mean field potential at the average interfacial site of residence for the prototropic moiety, F i s the Faraday constant, R is the universal gas constant and Tis the absolute temperature. In spite of the fact that eqn (1) is increasingly being used to rationalize acid-base behaviour at charged interfaces, very little quantitative work9* 1 9 7 23* 26 h as been done to isolate the principal factors contributing to the magnitude of the pK," component.Therefore, the present series of investigations were designed with the following major objectives : (a) to ascertain the magnitude of the non-electrostatic component to the values [i.e. the pK," value in eqn (l)] for a wide range of interfacially located 'simple' weak acids and bases and commonly used acid-base indicators; and (b) to determine the reasons for the difference, if any, between the pK," values of these interfacially located ' simple ' weak acids and bases and commonly used indicators and their p& values in bulk aqueous solution.In this first paper in the series, the acid-base equilibria of a number of 'simple' weak acids and weak bases in aqueous micellar solutions are examined and a quantitative assessment of their pK," values is given. There is an enormous pool of data in the literature concerned with the acid-base equilibria of 'simple' weak acids and bases in organic solvent-water mixtures. Some of these data are used to employ a procedure first developed by Fernandez and Fromherzg to compare the acid-base equilibria in organic solvent-water mixtures with those in aqueous non-ionic and charged micellar solutions. t Present address : CSIRO, Division of Chemicals and Polymers, G.P.O. Box 433 1, Melbourne, Victoria, 3001, Australia. 52 1522 ' Simple' Acid-Base Equilibria in Micelles pH-Titration results obtained in the present work for p-nitrophenol, p-(t-buty1)phenol and p-toluidine in both non-ionic and charged micellar systems are analysed. In addition, the results of other researchers for myristic ~tearylamine~~ and 4-octadecyloxy- 1 -naphthoic acid" in micellar systems are also examined.Since a number of biologically important molecules exhibit either weak acid or weak base behaviour there is an added incentive to investigate the acid-base equilibria of 'simple' weak acids and bases at lipid-water interfaces. For example, the apparent p& values of fatty bile acids,2s uncouplers of oxidative phosphorylation,' ph~sphatidylserine,'~~~~ phosphatidylethanolamine20 and some local anaesthetics15 located in rudimentary model biological membranes are all shifted relative to their p& values in pure water.Experiment a1 p-Nitrophenol, p-t-butylphenol and p-toluidine were purchased from Tokyo Kasei, Kogyo Co. The purity of each of these reagents was confirmed by established tests.29 Dodecyltrimethylammonium bromide (DTAB) was supplied by Tokyo Kasei and specially purified sodium dodecyl sulphate (SDS) by B.D.H. Chemicals Pty. Ltd. Both surfactants were recrystallized prior to use. Brij-3 5 and n-dodecyloctaoxyethylene glycol monoether (C12E,) were obtained from B.D.H. and Nikko Chemical Co., respectively. These surfactants were used as received. The inorganic reagents, tetraethylammonium chloride (TEAC), NaOH, HCl, KC1 and NaBr, were analytical grade and were employed without further purification.Buffer concentrates were from Merck. The aqueous solutions were prepared with Millipore-filtered water which had been distilled at least once before being filtered (conductivity < 1 x R-' cm-l and air-water surface tension equal to 72.0 mN m-' at 25 "C). All experiments were conducted at 25 "C. Both the conjugate acid and conjugate base forms of p-nitrophenol, p-(t-buty1)phenol and p-toluidine are relatively water soluble, hence high concentrations of surfactant were employed in an attempt to solubilize most of the species in the micellar phase. For the same reason micellar systems that had the same surface charge as the charged forms of the probe molecules were not examined. In order to ensure the removal of any undissolved weak acid or weak base, each solution was filtered before the pH-titration was performed.The solution under investigation was kept in a water-thermostatted reservoir fitted with a glass electrode and a double-junction reference calomel electrode. This dual electrode system was utilized so that the excessive drift in the pH-meter reading which is sometimes observed when a single combined electrode is used in concentrated micellar solution could be avoided.30 The upper compartment of the reference electrode contained a saturated KCl solution whilst the lower compartment contained a lop3 mol dm-3 solution of TEAC. The electrodes were standardized with a two-buffer adjustment procedure and the pH of the standardizing buffer solutions were always chosen to bracket the apparent p& of the weak acid or weak base in the system being examined. A Radiometer PHM84 research pH-meter was used to monitor the pH.The bulk pH of the solution in the reservoir was adjusted with small aliquots of concentrated solutions of NaOH or HCl in order to minimize dilution effects. After the solution had equilibrated, and a stable pH-meter reading had been reached, a sample was withdrawn from the reservoir and the u.v.-visible absorption spectrum was measured on a Varian Cary Model 210 spectrophotometer. This procedure was repeated until a full pH-titration profile was obtained. All the acid-base equilibria investigated were treated as being of the formC. J . Drummond, F. Grieser and T. W. Healy 523 where HX is the conjugate acid form of the molecule with charge z, H' is the proton and X is the conjugate base form of the molecule.The thermodynamic acid-base equilibrium constant for this reaction is given by The ratio of the concentrations of the conjugate acid-base forms could be determined from the spectra as a function of pH. In pure water the activity coefficient, y , of any singly charged species was approximated by the mean activity coefficient, y + , for HCl in water at the same total ionic strength.31 The influence of electrolyte on-the activity coefficient of a neutral species was assumed to be negligible, therefore the activity coefficient for a neutral species was set equal to unity. Since there is no available information on the mean activity coefficients of species residing within the interfacial microenvironment of micelles, the last term in eqn ( 3 ) had to be neglected for the apparent p& determinations in micellar solutions.The p& values in pure water and the apparent p& values in the micellar solutions were determined from the u.v.-visible absorption spectra as a function of the bulk aqueous pH using eqn (4) (4) Yx log - YH X A-AH, Ax - A m with a = where A is the absorbance at the long-wavelength band maximum of the conjugate base form, A,,,, at the particular pH being examined. AHx and A, are the absorbances at R,,, when the pH values are set such that only acid and base species are present, respectively. In order to ascertain whether or not the apparent p& of the molecules were well-defined, at least seven intermediate a values were looked at for each micellar solution.Examples of the change in the spectrum of p-nitrophenol, p-(t-buty1)phenol and p-toluidine with bulk aqueous pH can be found in ref. ( 3 2 ) . General Considerations The apparent pK, of the prototropic moiety of a weak acid (or weak base) residing within the aqueous interfacial microenvironment of a micelle can be split into a number of quantitative components. This section outlines the relationships and accompanying assumptions which allow one to derive these components. The initial part of this section closely follows reasoning presented el~ewhere.~~ 26 Nevertheless, it is restated herein because a knowledge of this reasoning is necessary for the understanding of the analysis and discussion of the results contained in this work and subsequent parts in this series.The apparent acid-base equilibrium for a weak acid (or weak base) located within an aqueous micellar interfacial phase can be represented as where subscripts i and w denote the interfacial phase and the bulk aqueous phase, respectively. The ' two-phase ' thermodynamic acid-base equilibrium constant for this reaction is defined by524 ' Simple' Acid-Base Equilibria in Micelles where a;, al,, and a;+ denote the activities of the various species in the phases. The experimentally determined apparent pKd has the form . . (8) Y'x Y L X and therefore = p c + log - . The standard Gibbs free energy of the 'two-phase' reaction given by eqn (9, AGO, and the ' two-phase ' thermodynamic acid-base equilibrium constant are related and with AGO expressed in terms of the standard chemical potentials for the species involved in the equilibrium where v/ is set equal convenient to define Also note that (9) to zero when a non-ionic interface is considered.In addition, it is 1 (Jpwf+p;-jp Fry 2.303 RT HX) - 2.303RT pKZ = Fv/ 2.303 RT (Jp;+ + p: - p;x). pPa = pK,"+ 1 2.303 RT - and the pK," component of eqn (1) and p E are related by PK," = p&+lOg-. Yfc Y h X The acid-base equilibrium for a weak acid (or weak base) in an organic solvent-water mixture (m) can be represented as HXg)$H;+Xg-l). (13) The ' single-phase ' thermodynamic acid-base equilibrium constant for this reaction is defined by pK," = -logTaR+ a: 'HX 1 If one assumes (a) that there are no specific molecular interactions which interfere with the interfacial acid-base equilibrium, and (b) that the solvent properties of an interfacial phase can be approximated by an organic solvent-water mixture with equivalent solvent properties, then for the particular equivalent organic solvent-water mixture and pK: and pK," differ solely by the work required to transfer the proton from the bulk aqueous phase to the interfacial phase, i.e.C. J .Drummond, F. Grieser and T. W. Healy The change of free energy for the transfer process is related to the medium effect on the proton, myH+, where H+ (std. state, w) -+ H+ (std. state, m) 525 by the formula33 /A;? -pLW+ = 2.303 RT log (18) PK; = pK," -log (19) and as a result The medium effect has also been referred to as the 'primary medium effect', the 'degenerate activity coefficient' and the 'distribution coefficient '.34 35 as to the nature of the medium effect on the proton.The problem is that a direct measurement of the medium effect on a single ionic species is not possible and it can only be estimated by making various non- thermodynamic assumptions. In this series of investigations the procedure of Fernandez and Fromherzg has been followed and it has been assumed that for a particular organic solvent-water mixture the log ,y,+ value can be approximated by the log ,y, value for HCl in the same organic solvent-water mixture. The justification for taking this approach will be given in the forthcoming discussion. The values of logmy, for HC1 in the various organic solvent-water mixtures were obtained from the change in the standard potential, E O , for the cell Pt/H,(g), HCl in solvent s, AgCl/Ag as an organic solvent is added to pure water.The relationship is given by3' At present, there is considerable - (,E" - ,E") F 2RTln 10 logmy, = where ,E" and ,E" are the standard potentials for the cell when pure water and an organic solvent-water mixture are the solvents, respectively. The value used for ,E" was 0.22234 V at 25 0C.37 The values for ,E" were either obtained from the compilation of Feakins and French38 or the one of Bates.34 The one exception to this was the ,E" value for the 82 weight % 1,4-dioxane-water mixture which was taken from the work of Danyluk et af.39 It should be mentioned that all the standard potentials used were based on the molar scale and as a result the calculated my, values relate to molarity.Fig. 1 shows logmy+ for HCl as a function of the dielectric constant of 1,4-dioxane-water, ethanol-water and methanol-water mixtures. The values for the dielectric constants were obtained by interpolation of the data of AkerloP' and Critchfield et aL4' For a particular weak acid (or weak base) the p& values obtained in micellar solutions and in organic solvent-water mixtures are, in general, referred to the pEi, value determined in pure water. Therefore it is convenient to define ApK," may be separated into an electrostatic and a non-electrostatic component :42 A P C = (APK,")~, +(APK,")none,* (24) The electrostatic component is due to the different amounts of work required to charge526 2.0 i-* E 4 8 1.0 ' Simple ' Acid-Base Equilibria in Micelles 20 40 60 dielecnic constant Fig.1. logmy, Values as a function of the dielectric constant of methanol-water (e), ethanol-water (0) and 1,4-dioxane-water (0) mixtures. the ions in two media of different dielectric constant. It can be estimated from the Born equation :42* 43 where N is Avogadro's number, e is the elementary charge, E is the dielectric constant and rj is the ionic radius of species j . The non-electrostatic component incorporates all the additional influences on A p e , such as specific solute-solvent interactions. Using the relationships given in eqn (19) and (24), eqn (22) can be rewritten as: (26) AP& = (ApK,">el + (APK,")nont?l - log rnYHf. Hence, if the quantity (ApC),,,,, -log myH+ is either small in magnitude compared with (ApK,"),, or equivalent for different solvent mixtures with the same dielectric constant, one would expect to observe for a particular molecule the same ApFi behaviour as a function of solvent dielectric constant irrespective of the particular kind of solvent mixture.In this and subsequent parts in this series, unless specifically indicated to the contrary, it is taken for granted that in charged micelles the prototropic moieties of the guest molecules reside, on average, in the plane of the surface charge. As will become apparent the results justify this assumption. Independent n.m.r. also indicate that the vast majority of non-lipoidal micelle solubilized aromatic molecules reside on average within the interfacial headgroup region of charged micelles.Results and Discussion The logmy, data in fig. 1 were utilized to construct reference ApK: curves for the molecules investigated from their pK," values in organic solvent-water mixtures of1.0 .- m 3 0.5 C. J . Drummond, F. Grieser and T. W. Healy I I 1 527 2 0 4 0 6 0 dielectric constant Fig. 2. p-Nitrophenol ApKj reference values as a function of the dielectric constant of methan~l--water~~~~~-~~ (a) and ethano1-waterg9 (0) mixtures. known dielectric constant. Fig. 2-7 show the ApKi values as a function of the solvent dielectric constant. All the pK," values employed in this study were obtained from the literature. Some of the pK," values were based on the molality scale so these were converted into molarity units before subtracting log The various sources of the pK," values43.45-61 are indicated in the relevant figure captions. The ApKA curve for p-(t-butyl)phenol, fig. 3, was actually constructed from ApK," values for phenol. This was done because literature pK," data for phenol are far more abundant. From the work of Parsons and Roche~ter'~ with substituted phenols, it is clear that the ApK," behaviour of phenol and p-(t-buty1)phenol are equivalent. For myristic acid and stearylamine, fig. 5 and 6, the ApK," behaviour with dielectric constant was assumed to be the same as for their water soluble homologues propionic acid and hexylamine. Similarly, 4-octadecyloxy- 1 -naphthoic acid, fig. 7, was assumed to have the same ApK," behaviour as benzoic acid. As shown in fig. 2, there is a small difference between the methanol-water and ethanol-water ApKi values of p-nitrophenol as a function of solvent dielectric constant.Fig. 3 indicates that the ApKi behaviour of phenol with solvent dielectric constant is equivalent for dioxane-water and ethanol-water mixtures. Interestingly, however, there is a significant difference between this ApKA behaviour and that found for meth- anol-water mixtures. Evidently, there is some kind of aberrant specific solute-solvent interaction in this particular methanol-water system. The reference ApK: values for p - toluidine as a function of dielectric constant, fig. 4, also show some solvent specificity as indicated by the different dioxane-water and ethanol-water results. In the organic solvent-water mixtures examined for propionic acid, hexylamine and benzoic acid, ApKL is a unique function of solvent dielectric constant.This can be seen in fig. 5-7. In this study, it has been assumed that the log (yk/yLx) term in eqn (1 2) is negligibly small. It has also been assumed that there is no contribution to pK," values due to specific528 ' Simple ' Acid-Base Equilibria in Micelles 2.0 .- * 3 1 .o 20 40 60 dielectric constant Fig. 3. p-(t-Buty1)phenol ApKl reference values, based on phenol, as a function of the dielectric constant of rnethan~l--water~~*~~ (a), e t h a n o l - ~ a t e r ~ ~ ~ ~ ~ (0) and 1 ,4-dioxane-water51 (0) mixtures. 2 0 40 60 dielectric constant Fig. 4. p-Toluidine ApK: reference values as a function of the dielectric constant of ethanol- ~ a t e r ~ ~ , ~ ~ , ~ ~ (0) and 1,4-dio~ane-water~~ (0) mixtures.C.J . Drummond, F. Grieser and T. W. Healy 529 20 40 60 dielectric constant Fig. 5. Myristic acid ApKi reference values, based on propionic acid, as a function of the dielectric constant of methan~l-water~~.~~ (O), e t h a n ~ l - w a t e r ~ ~ * ~ ~ . ~ ’ (0) and 1,4-dioxane-~ater~~-~~ (0) mixtures. -1.0 -2.0 2 0 40 6 0 dielectric constant Fig. 6. Stearylamine ApKi reference values, based on hexylamine, as a function of the dielectric constant of methanol-wateP (0) and ethan~l-water~~ (0) mixtures.530 ' Simple' Acid-Base Equilibria in Micelles fl 20 40 60 dielectric constant Fig. 7. 4-Octadecyloxy-1-naphthoic acid ApK: reference values, based on benzoic acid, as a mixtures. function of the dielectric constant of rnethan~l-water~~ (a) and ethan~l-water~~,~~.~' (0) molecular interactions. Thus by comparing a ApK," value with the plot of reference ApP; values as a function of the solvent dielectric constant, one can estimate the effective dielectric constant of the interfacial microenvironment of micelles.As the ApKA values of myristic acid, stearylamine and 4-octadecyloxy- 1 -naphthoic acid apparently respond uniquely to changes in the solvent dielectric constant, these molecules can provide unambiguous estimates of the interfacial Eeff at their average site of residence. For p- nitrophenol, p-(t-buty1)phenol and p-toluidine there will obviously be some ambiguity associated with the Eeff estimates. Tables 1, 2 and 3 contain the pH-titration results obtained for p-nitrophenol, p-(t- buty1)phenol and p-toluidine, respectively.Table 4 contains the pH-titration results of Ptak et aL2' for myristic acid and stearylamine and those of Lovelock et aI.l7 for 4- octadecyloxy-1-naphthoic acid. Also included in all the tables are the ApK," and Eeff estimates. Where the reference ApKA behaviour was not a unique function of solvent dielectric constant, eeff values referring to each calibrating organic solvent-water mixture have been given. For the charged micellar systems the pK," values were determined from the known micellar surface potential^,^^^^^ the P K , " ~ ~ values and eqn (1). The electrostatic surface potential of a DTAB micelle in a 2, 5 and 10 weight % DTAB solution is considered to be + 114, +98 and +91 mV, respectively, whilst in the presence of 4 mol dm-3 NaBr it is considered to be + 18mV.19924 The surface potential of a CTAB micelle in a 0.05 mol dm-3 CTAB solution is considered to be + 141mV.19 With water soluble molecules the possibility exists that the PK,"~~ value is a composite value, comprising contributions from species within the interfacial phase and the bulk aqueous phase.In this study we attempted to avoid this occurrence by using highC. J. Drummond, F. Grieser and T. W. Healy 53 1 Table 1. pH-Titration results for p-nitrophenol in pure water and aqueous micellar solutions with the corresponding ApK," and ceff estimates 'max/nm solution WtYo" ROH (RO-) pKibS APK," Eef? water 0 3 14 (397) 7.15 f 0.05 0 - Brij-35 5 312 (396) 7.86k0.12 0.71 3 8 f 3 (M) 3 7 f 4 (E) Brij-35 10 3 10 (396) 8.08 & 0.17 0.93 35 f 2 (M) 31 f 4 (E) DTAB 2 3 16 (399) 6.27 k 0.06 1 .05c 34 f 1 (M) < 31 (E) DTAB 5 316 (398) 6.45f0.06 0.96' 35f 1 (M) < 31 (E) DTAB 10 315 (397) 6.62k0.03 1.01" 35f 1 (M) < 31(E) DTAB-NaBr (4 mol dm-3) 5 314 (393) 8.01 k0.03 1.16' < 34 (M) <31 (E) a Weight YO surfactant in solution.M or E denotes whether methanol-water or ethanol-water mixtures were used as the reference, respectively. ' Calculated with eqn (1) and the 'yo value estimated by utilizing the ET(30) molecule. l 9 s Z 4 Table 2. pH-Titration results for p-(t-buty1)phenol in pure water and aqueous micellar solutions with the corresponding ApK,O and cerf estimates water 0 274 (291) 10.31 f 0.02 0 - CI2El-4 5 277 (293) 11.96 f 0.05 1.65 33f 1 C12El3 10 276 (293) 12.04 & 0.03 1.73 3 1 f l DTAB 2 276 (296) 9.83 0.02 1 .45c 37f 1 DTAB 5 277 (297) 10.07 f 0.02 1.42' 37+ 1 DTAB 10 275 (296) 10.30 f 0.03 1.53' 35f 1 DTAB-NaBr (4 mol dm-3) 5 277 (297) d - - a Weight YO surfactant in solution.1,4-dioxane-water and ethanol-water reference ApK: curve used. ceff values based on methanol-water all < 34. ' Calculated with eqn (1) and the ry, value estimated by using the ET(30) molecule.1g~24 pK:hs was not well defined varying from 11.45 at a = 0.092 to 10.75 at a = 0.925. concentrations of surfactants. Nevertheless, the PK,"~" results for p-(t-buty1)phenol in 5 wt % DTAB solution with 4 mol dm-3 NaBr and those for p-toluidine in both C,,E, and SDS solution suggest that a high percentage of the species in these systems may not have partitioned into the interfacial phase.Alternatively, the finding that the pK,"bs value of p-toluidine in micellar SDS solution is not well-defined, table 3, can be accounted for if there is specific interaction between the positively charged p-toluidinium ion and a negatively charged sulphate headgroup. Indeed there is some other evidence6,. 63 which suggests that this may be the case. Henceforth the p-toluidine results will be ignored in the discussion. It is clear from the results of tables 1, 2 and 4 that the ApKZ values of p-nitrophenol, p-(t-butyl)phenol, myristic acid, stearylamine and 4-octadecyloxy- 1 -naphthoic acid in532 ' Simple' Acid-Base Equilibria in Micelles Table 3. pH-Titration results for p-toluidine in pure water and aqueous micellar solutions with the corresponding ApK," and eeff estimates 4n,x/nm solution wt Yo" RNH, pK;" APK," &elfb - water 0 285 5.05 f 0.03 0 C12E8 5 288 4.62 + 0.05 - 0.43 59 f 2 (D) 65 k 2 (E) C12E8 10 289 4.39 + 0.02 -0.66 51 f 1 (D) 59f 1 (E) SDS 2 286 6.60 -, 6.89" - - SDS 5 286 6.63 +6.96" - - SDS 10 286 6.28-6.56' - - a Weight YO surfactant in solution.D or E denotes whether 1,4-dioxane-water or ethanol-water mixtures were used as the reference, respectively. ' pKibS not well defined; low a value -, high a value. Table 4. Literature apparent p q results for myristic acid,27 stearylamineZ7 and 4-octadecyloxy- 1 - naphthoic acid17 in aqueous micellar solution with the corresponding ApK: and eePf estimates from this study molecule PK," a pKibS APK: 'eff myristic acid 5.0 6.7b + 1.7 31 f 4 stearylamine 10.7 8.95b - 1.75 40k4 4-octadecyloxy- 1 -naphthoic acid 4.2 6.60" + 2.4 35+ 1 4-octadecyloxy- 1 -naphthoic acid 4.2 4.20d + 2.38" 35f 1 ~~ a Estimated p& value in water based on water soluble homologues.Micellar Triton X-100 solution. 0.005 mol dm-3 Brij-35 or 0.05 mol dm-3 CI2E8 solution. 0.05 mol dm-3 CTAB solution. " Calculated using eqn (1) and the yo value estimated by utilising the &(30 rn01ecule.'~ non-ionic micelles comprising surfactant molecules with poly(ethy1ene oxide) head- groups simply reflect the interfacial solvent properties. These interfacial solvent properties are well characterized by an effective dielectric constant of 35 + 5. Similarly, the calculated ApK: values for p-nitrophenol, p-(t-buty1)phenol and 4- octadecyloxy- 1 -naphthoic acid in cationic micelles comprising surfactant molecules with quaternary ammonium headgroups, are indicative of an interfacial microenvironment which possesses an effective dielectric constant of ca.35. This finding also validates the practice of splitting PK,"~" values into an electrostatic surface potential component and a non-electrostatic component as is done in eqn (1). Moreover it clearly establishes the independence of the two components. Using an identical procedure to the one employed in this work, Fernandez and Fromherzg obtained an eeff value of 32+ 1 for the interfacial microenvironment of non- ionic Triton X-100 micelles. This value was determined by comparing the ApK," values of 7-hydroxy-4-undecylcoumarin and 7-amino-4-heptadecylcoumarin in Triton X- 100 micelles with the Ap& behaviour of their water soluble homologues in 1,4- dioxane-water mixtures.The eeff values gauged in the present investigation are also in agreement with a number of other estimates gained by utilizing totally different techniques. From the solvent-sensitive u.v.-absorption spectrum of solubilized dode- cylpyridinium iodide, Mukerjee et aL6* estimated the Eeff value for the interfacial regionC. J. Drummond, F. Grieser and T. W. Healy 533 of Brij-35 micelles to be 36. This estimate was based on alcohols and alcohol-water mixtures as the reference solvents. Law65 employed the solvent-sensitive fluorescence emission maximum of 2-[6-(2,2-dicyanovinyl)-3,4-dihydro-2,2,4-trimethyl- 1 (2H)-qui- noyllethylbenzoate, with a range of neat organic solvents and alcohol-water mixtures as reference solvents, and estimated that the eeff values for the interfacial micro- environments of Triton X-100 and CTAB micelles are 28 & 8 and 36 & 8, respectively.Utilizing the solvatochromic visible absorption-band maximum of ET(30) and n-alcohols, ethanol-water mixtures and dioxane-water mixtures as reference solvents, both Zachariasse et al." and Drummond et al." have obtained Eeff estimates of 29-36, 28-33 and 27-30 for DTAB, CTAB and C,,E, micelles, respectively. The significance of the E~~~ value characterizing the interfacial region of non-ionic micelles comprising surfactant molecules with poly(ethy1ene oxide) headgroups may be ascertained by comparing the interfacial solvent properties with those of poly(ethy1ene glycol)-water systems.Static dielectric constant measurements of poly(ethy1ene glycoljwater mixtures indicate that bulk solution dielectric constants of 30-40 are attained at an average of 0.5-1 .O water molecules per poly(ethy1ene oxide) subunit." Thus the Eeff values are consistent with the interfacial region of the non-ionic micelles possessing low water activity. It seems likely that the low Eeif values found for the interfacial microenvironment of CTAB and DTAB micelles may also be attributed to a low interfacial water activity.ss~s8 Nevertheless it should be mentioned that the low interfacial E~~~ values may also be a result of the hydrogen-bond donor properties of the water in the interfacial regions being different from that of bulk water, and/or the presence of electrostatic image interactions caused by the proximity of the low dielectric hydrocarbon core.69 Conclusions The acid-base equilibria of a wide range of ' simple ' weak acids and bases residing within the interfacial microenvironment of both non-ionic and charged micelles have been investigated. Concordant estimates have been obtained for the effective dielectric constant of the interfacial region of both non-ionic micelles comprised of surfactant molecules with poly(ethy1ene oxide) headgroups and charged micelles composed of surfactants with quaternary ammonium headgroups. The concordance of the Eeff estimates determined for the same interfacial microenvironment from the acid-base behaviour of the different weak acids and bases indicates that (a) for the molecules and systems examined, with the probable exception of p-toluidine in micellar SDS solution, the differences between the p c and pK," values can primarily be ascribed to the difference between the mean solvent properties of the interfacial phase and the bulk aqueous phase ; (b) interfacial solvent effects on the acid-base equilibria investigated can be likened to those of an organic solvent-water mixture; (c) for the molecules and systems studied, with the probable exception of the p-toluidinium ion in SDS micelles, there is no notable contribution to pK," from any kind of specific molecular interaction within the interfacial region; ( d ) the medium effect on the proton, myH+, can reasonably be approximated by the mean ionic medium effect on HCl, J + ; and (e) for the weak acids and weak bases investigated in this work the logy!JyL, term in eqn (12) is negligibly small.This work was supported by the Australian Research Grants Scheme. References 1 G. S. Hartley and J. W. Roe, Trans. Faraday SOC., 1940, 36, 101. 2 P. Mukerjee and K. Banerjee, J. Phys. Chem., 1964, 68, 3567. 3 S. H. Yalkowsky and G. Zografi, J . Colloid Interface Sci., 1970, 34, 525.534 ' Simple' Acid-Base Equilibria in Micelles 4 F. Tokiwa, A h . Colloid Interface Sci., 1972, 3, 389. 5 M. Montal and C. Gitler, J. Bioenerg., 1973, 4, 363. 6 P. Fromherz and B. Masters, Biochim. Biophys. Acta, 1974, 356, 270. 7 E. P. Bakker, J. C. Arents, J.P. M. Hoebe and H. Terada, Biochim. Biophys. Acta, 1975, 387, 491. 8 N. Funasaki, J. Colloid Interface Sci., 1977, 60, 54. 9 M. S. Fernandez and P. Fromherz, J. Phys. Chem., 1977, 81, 1755. 10 W. L. C. Vaz, A. Nicksch and F. Jahnig, Eur. J. Biochem., 1978, 83, 299. 11 T. Mashimo, I. Ueda, D. D. Shieh, H. Kamaya and H. Eyring, Proc. Natl Acad. Sci. USA, 1979, 76, 12 J. Frahm, S. Diekmann and A. Haase, Ber, Bunsenges. Phys. Chem., 1980, 84, 566. 13 G. Cevc, A. Watts and D. Marsh, Biochemistry, 1981, 20, 4955. 14 S. Lukac, J. Phys. Chem., 1983, 87, 5045. 15 J. Garcia-Soto and M. S. Fernandez, Biochim. Biophys. Acta, 1983, 731, 275. 16 P. Fromherz and R. Kotulla, Ber. Bunsenges. Phys. Chem., 1984, 88, 1106. 17 B. Lovelock, F. Grieser and T. W. Healy, J. Phys. Chem., 1985, 89, 501.18 B. Lovelock, F. Grieser and T. W. Healy, Langmuir, 1986, 2, 443. 19 C. J. Drummond, F. Grieser and T. W. Healy, Faraday Discuss. Chem. Soc., 1986, 81, 95. 20 F. C. Tsui, D. M. Ojcius and W. L. Hubbell, Biophys. J., 1986, 49, 459. 21 G. V. Hartland, F. Grieser and L. R. White, J. Chem. SOC., Faraday Trans. I , 1987, 83, 591. 22 C. J. Drummond and F. Grieser, Photochem. Photobiol., 1987, 45, 19. 23 C. J. Drummond and F. Grieser, Langmuir, 1987, 3, 855. 24 C. J. Drummond, F. Grieser and T. W. Healy, Chem. Phys. Lett., 1987, 140, 493. 25 J. Kibblewhite, C. J. Drummond, F. Grieser and T. W. Healy, J. Phys. Chem., 1987, 91, 4658. 26 C. J. Drummond, F. Grieser and T. W. Healy, J. Phys. Chem., 1988, 92, 2604. 27 M. Ptak, M. Egret-Charlier, A. Sanson and 0.Bouloussa, Biochim. Biophys. Acta, 1980, 600, 387. 28 A. F. Hofmann and A. Roda, J. Lipid Res., 1984, 25, 1477. 29 G. Ackermann, L. Sommer and W. I. Stephen, Pure Appl. Chem., 1985, 57, 845. 30 A. L. Underwood, Anal. Chim. Acta, 1977, 93, 267. 3 1 H. S. Harned and B. B. Owen, The Physical Chemistry of Electrolyte Solutions (Reinhold, New York, 32 C. J. Drummond, Doctoral Dissertation (The University of Melbourne, 1987). 33 R. G. Bates, Determination ofpH: Theory and Practice (Wiley, New York, 2nd edn, 1973), p. 213. 34 R. G. Bates, in Hydrogen-bonded Solvent Systems, ed. A. K. Covington and P. Jones (Taylor and 35 0. Popovych, Crit. Rev. Anal. Chem., 1970, 1, 73. 36 R. G. Bates, Determination of pH: Theory and Practice (Wiley, New York, 2nd edn, 1973), p. 270.37 R. G. Bates, Determination of pH: Theory and Practice (Wiley, New York, 2nd edn, 1973), p. 334. 38 D. Feakins and C. M. French, J. Chem. SOC., 1957, 2581. 39 S. S. Danyluk, H. Taniguchi and G. J. Janz, J. Phys. Chem., 1957, 61, 1679. 40 G. Akerlof, J. Am. Chem. SOC., 1932, 54,4125. 41 F. W. Critchfield, J. A. Gibson and J. L. Hall, J. Am. Chem. SOC., 1953,75, 1991. 42 R. G. Bates and R. A. Robinson, in Chemical Physics of Ionic Solutions, ed. B. E. Conway and R. G. 43 G. H. Parsons and C. H. Rochester, J. Chem. SOC., Faraday Trans. 1, 1974, 70, 1058. 44 K. N. Ganesh, P. Mitra and D. Balasubramanian, J. Phys. Chem., 1982, 86, 4291, and references 45 C. L. De Ligny, H. Loriaux and A. Ruiter, Recueil, 1960, 80, 725. 46 R. A. Robinson and R. G. Bates, J. Res. Nut. Bur. Stand., Sect. A , 1966, 70, 553. 47 R. Gaboriaud, Ann. Chim. (Paris), 1967, 2, 201. 48 J. Juillard, C.R. Acad. Sci., Ser. C, 1969, 268, 2251. 49 R. Thuaire, J. Chim. Phys., 1972, 69, 23. 50 C. C. Panichajakul and E. M. Woolley, Anal. Chem., 1975, 47, 1860. 51 J. M. Sanchez-Ruiz, J. Llor and M. Cortijo, J. Chem. SOC., Perkin Trans. 2, 1984, 2047. 52 B. Gutbezahl and E. Grunwald, J. Am. Chem. SOC., 1953, 75, 559. 53 W. J. Gelsema, C. L. De Ligny and G. F. Visserman, Recueil, 1965, 84, 1129. 54 H. P. Marshall and E. Grunwald, J. Am. Chem. SOC., 1954, 76, 2000. 55 A. L. Bacarella, E. Grunwald, H. P. Marshall and E. L. Purlee, J. Org. Chem., 1955, 20, 747. 56 G. Papanastasiou, G. Stalidis and D. Jannakoudakis, Bull. SOC. Chim. Fr., 1984, 9-10, 255. 57 E. Grunwald and B. J. Berkowitz, J. Am. Chem. SOC., 1951, 73, 4939. 58 H. S. Harned and T. R. Dedell, J. Am. Chem. SOC., 1941, 63, 3308. 59 L. G. Van Uitert and G. G. Haas, J. Am. Chem. SOC., 1953, 75,451. 60 0. Dusart, C. Piekarski and S. Piekarski, J. Chim. Phys., 1975, 72, 97. 61 C. L. De Ligny, Recueil, 1960, 79, 731. 62 T. Yamashita, H. Yano, S. Harada and T. Yasunaga, J. Phys. Chem., 1984, 88, 2671. 51 14. 3rd edn, 1958), p. 716. Francis, London, 1968), p. 49. Barradas (Wiley, New York, 1964), p. 211. therein.C. J . Drummond, F. Grieser and T. W. Healy 535 63 T. Yamashita, M. Sumino, H. Yano, S. Harada and T. Yasunaga, Bull. Chem. SOC. Jpn, 1984, 57, 2352. 64 P. .Mukerjee, J. R. Cardinal and N. R. Desai, in Micellization, Solubilization and Microemulsions, ed. 65 K. Y. Law, Photochem. Photobiol., 1981, 33, 799. 66 K. A. Zachariasse, N. Van Phuc and B. Kozankiewicz, J. Phys. Chem., 1981, 85, 2676. 67 D. Dunstan, Dept. Physical Chemistry, The University of Melbourne, personal communication. 68 M. Gutman, Methods Biochem. Anal., 1984, 30, 1. 69 C. Ramachandran, R. A. Pyter and P. Mukerjee, J . Phys. Chem., 1982, 86, 3198. K. L. Mittal (Plenum Press, New York, 1977), vol. 1, p. 241. Paper 7/00 103G ; Receiiied 2 1st December, 1987