J . Phys. Chem. 1989, 93, 4219-4226
4219
Speclflc Counterion Effects on Indicator Equilibria in Micellar Solutions of Decyl Phosphate and Lauryl Sulfate Surfactants Zhen-Min He, Patrick J. O'Connor, Laurence S. Romsted,* Department of Chemistry, Rutgers, The State University of New Jersey, New Brunswick, New Jersey 08903
and Din0 Zanette* Departamento de Quimica, Universidade Federal de Santa Catarina, 88.000Florianopolis, Santa Catarina. B r a d (Received: August 16, 1988; In Final Form: December 20, 1988)
The pseudophase ion exchange (PIE) model provides both a qualitative and quantitative interpretation of aqueous micellar effects on reaction rates and equilibria for a variety of thermal reactions between organic molecules and ions in the presence and absence of buffer. Our results show that the PIE model is also applicable to alkali-metal salts of decyl phosphate monoanion micelles, MDP (M = Na, K, Rb, and Cs). We measured the ratio of acid to base forms of the spectrophotometric indicator pyridine-2-azo-p-dimethylaniline, PADA, as a function of added salt MCI, up to 0.4 M in 0.08 M MDP, pH range 4-6, succinate buffer at 50 O C . Similar experiments were run in micellar solutions of sodium lauryl sulfate, NaLS, with added MCI under the same conditions to check for special effects caused by the phosphate monoanion head group. None were found. The assumptions of the PIE model are well obeyed provided we use the measured pH to calculate the activity of the proton at all salt concentrations and express the quantity of counterions in the aqueous phase as activities and not concentration units. Precise results also require correcting for specific salt effects on the response of the pH electrode. Our estimated value of pKAv,the acidity constant of micellar-bound PADA, is only slightly smaller than its value in water in both MDP and NaLS micelles, consistent with the commonly held view that reaction occurs at the water-rich micelle surface and not in the hydrocarbon core. The relative values of the ion-exchange constants between the alkali-metal cations and the proton, KHM,follows a Hofmeister series in both surfactants, Cs > Rb > K > Na, Le., decreasing selectivity with increasing size of the hydrated cation. The alkali-metal cation selectivity order for several different phospholipid vesicles follows the opposite trend, suggesting that specific counterion interactions in micelles and vesicles with phosphate head groups are fundamentally different.
Introduction Ionic surfactants in solution form a wide variety of dynamic aggregates, sometimes called association colloids, including micelles, monolayers, microemulsions, inverse micelles, vesicles, and biological m e m b r a n e ~ . l - ~A number of the properties of these aggregates, such as the critical concentration for aggregate formation, aggregate size and shape, aggregate and phase stability, and the binding of ions and molecules, depend on both the type and concentration of c o ~ n t e r i o n s . l - Many ~ association colloids alter the rates of chemical reactions and shift the acidity constants of pH indicators.e22 Bimolecular reactions and equilibria also (I) Fendler, J. Membrane Mimetic Chemistry; Wiley-Interscience: New York, 1982. (2) Surfactants in Solution; Lindman, B., Mittal, K . L., Eds.; Plenum Press: New York, 1984; Vol. 1-3. (3) Modern Trends of Colloid Science in Chemistry and Biology; Eicke, H.-F., Ed.; Birkhauser Verlag: Basel, 1985. (4) Mackay, R. A. Adv. Colloid Interface Sci. 1981, 15, 131. (5) Fernandez, M. S.; Fromherz, P. J. Phys. Chem. 1977, 81, 1755. ( 6 ) Romsted, L. S. In Surfuctants in Solution; Mittal, K . L., Lindman, B., Eds.: Plenum Press: New York, 1984; Vol. 2, p 1015. (7) Drummond, C. J.; Grieser, F.; Healy, T. W. J . Phys. Chem. 1988, 92, 2604. (8) Fuhrhop, J.-H.; Penzlin, G.; Tank, H. Chem. Phys. Lipids 1987, 43, 147. (9) Warr, G . G.; Evans, D. F. Langmuir 1988, 4, 217. ( I O ) Sudholter, E. J. R.; Van de Langkruis, G . B.: Engberts, J. 8. F. N . R e d . Trav. Chim. Pays-Bas 1980, 99, 73. ( 1 I) Chaimovich, H.; Aleixo, R. M. V.; Cuccovia, I. M.; Zanette, D.; Quina, F. M . In Solution Behavior of Surfactants: Theoretical and Applied Aspects: Mittal, K . L., Fendler, E. J., Eds.; Plenum Press: New York, 1982; VOl. 2, p 949. (12) Mukerjee, P.: Banerjee, K. J . Phys. Chem. 1964, 68, 3567. (I3) Martinek, K.; Yatsimirski, A. K.; Levashov, A. V.; Berezin. I. V. I n Micellization, Solubilization and Microemulsions: Mittal, K . L., Ed.; Plenum Press: New York, 1977; Vol. 2, p 489. ( 14) Romsted, L. S. In Micelliration, Solubilization and Microemulsions; Mittal, K . L., Ed.; Plenum Press: New York, 1977: Vol. 2, p 509. ( I S ) Quina, F. H.; Chaimovich, H . J . Phys. Chem. 1979, 83, 1844. (16) Gaboriaud, R.: Charbit, G.; Dorion, F. In Surfactants in Solution; Mittal, K. L., Lindman. B., Eds.; Plenum Press: New York, 1984; Vol. 2, p 1191.
0022-3654/89/2093-4219$01.50/0
show a marked dependence on counterion type and concentration.5,6,13-15s20-23Thermodynamically stable and structurally "simple" aqueous micelles are often used as model systems for the far more complex and metastable biological membranes because their interfaces are structurally similar. Bimolecular reactions and equilibria respond in fundamentally the same way to changes in surfactant structure such as alkyl chain length, head-group charge, and counterion type as well as changes in surfactant and salt concentration. A number of quantitative models have been developed over the past several decades for interpreting these result^.^-^^"-^^ The pseudophase ion exchange (PIE) model provides a quantitative description of aqueous micellar effects on both reaction rates and indicator equilibria, and its strengths and weaknesses have been reviewed repeatedly.'*6~'o~'1~20-22 The basic assumptions of the PIE model are illustrated in Figure l : (a) Micelles act as a separate phase, uniformly distributed throughout the solution, and the medium properties of this phase are independent of solution composition. Thus, the ionization of a micellar-bound indicator is described by an intrinsic acidity constant, KAm, for ionization within the micellar pseudophase (AH,,, + A, + H,). (b) The fraction of bound counterions, 0.is constant and independent of surfactant concentration, counterion concentration, and counterion type. (c) The micelle surface acts like an ion-exchange resin, and the (17) Lelievry, J.; Gaboriaud, R. J . Chem. Soc., Faraday Trans. I 1985, 81, 335. ( I 8) Bunton, C. A,: Moffatt, J . R. In Colloids and Surfactants: Fundamentals and Applications; Barni, E., Pelizzetti, E., Eds.; Reprinted from Ann. di Chim. 1987, 77, 1 17 and references therein. ( 1 9) Drummond, C. J.: Grieser, F.; Healy, T. W. Faraday Discuss. Chem. Soc. 1986, 81, 95. (20) Bunton, C. A,; Romsted, L. S. In The Chemistry of the Functional Groups. Supplement B The Chemistry of Acid Derioarioes; Patai, S . , Ed.; Wiley-Interscience: London, 1979; Part 2, p 945. (21) Romsted, L. S. J . Phys. Chem. 1985, 89, 5107, 51 13. (22) Bunton, C. A,; Savelli, G . Ado. Phys. Org. Chem. 1986, 22, 213. (23) Bunton, C. A. I n Reaction Kinetics in Micelles; Cordes, E. H..Ed.; Plenum Press: New York, 1973; p 73.
1989 American Chemical Society
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The Journal of Physical Chemistry, Vol. 93, No, 10, 1989
H e et ai. monomer in water are the same. A substantial amount of work on the alkyl phosphate surfactants including preparation32and micellar properties such as critical micelle concentration (cmc) and Krafft point of alkyl phosphate m ~ n o a n i o n ~ ~ and d i a n i ~ nsurfactants ~~ of varying chain lengths. Recently, spectroscopic methods, including ESR and 3'P and I3C N M R , were used to estimate micelle size, shape, degree of counterion binding and internal chain mobility.35 Measurements were also made in micellar solutions of sodium lauryl sulfate, NaLS, for comparison because, unlike the phosphate group, the sulfate group is a strong acid and cannot contribute intermolecular hydrogen bonds. In an earlier paper we measured, spectrophotometrically, the effect of added NaCl to aqueous micelles of sodium decyl phosphate monoanion, NaDP, and NaLS on the ratio of the conjugate base to acid forms of a completely micellar-bound indicator, pyridine-2-azo-p-dimethylaniline, PADA.24 One of the most important findings of that study is that when the quantity of protons in the aqueous phase is measured by using a pH meter, the quantities of protons and sodium ions in the aqueous phase must be expressed as activities instead of concentrations to obtain a good correlation. Here we compare the effect of up to 0.4 M added MCI (M = K, Rb, and Cs) on the indicator ratio of PADA in micellar solutions of 0.08 M decyl phosphate monoanion, 0.02 M succinate buffer a t 50 OC. To prevent mixing of counterions, both the surfactant and buffer were prepared by neutralization with a metal hydroxide having the same alkali-metal ion as the added salt. Experiments were also carried out in micellar solutions of 0.04 M NaLS and 0.02 M sodium succinate as buffer a t 50 "C contain the same set of alkali-metal salts. However, because all the solutions contained sodium ions in addition to one of the other metal ions, the interpretation is less rigorous. Many of the experimental procedures and control experiments are described In excruciating detail in the previous paper24 and will not be repeated here. IS published
0-Lon 0
Figure 1. Cartoon of a partial cross section of the region near the in-
terface of a MDP micelle showing the micellar and aqueous pseudophases separated by the Stern layer. The Stern layer contains phosphate head groups with covalently bound protons, alkali-metal ion, and proton counterions and the two forms of the indicator, AH+ and A. The aqueous phase contains MDP monomer, protons, metal ions, and the various forms of succinate buffer, BHz, BH-, and B2-. distribution of the proton, H, and an alkali-metal ion, M, is described by an ion-exchange constant, KHM.In these experiments, succinate buffer is added to control the p H of the aqueous phase. Because succinate buffer is very water soluble and its mono- and dianionic forms are of like charge to the micelles, we assume that the buffer species are only in the aqueous phase and do not buffer the micelle ~ u r f a c e . ~ ~ , ~ ~ T o date, most of the published work on specific salt effects on indicator equilibria in micelles is in alkaline solutions of cationic surfactants, and only a small fraction is on specific salt effects on reaction rates and indicator equilibria in anionic surfact a n t ~ . ' . ~ . ~ We ~ . selected ~ ~ ~ ~ the ~ . alkali-metal ~ ~ - ~ ~ salts of decyl phosphate monoanion. MDP, because the micellar interface is a simple model for anionic membrane phospholipids such as phosphatidic acid, PA, phosphatidyl glycerol, PG, or phosphatidyl serine, PS. To ensure that the head groups were primarily in their monoanion form and that the surface charge density was not affected by the presence of significant amounts of phosphoric acid monodecyl ester a t low pH or decyl phosphate dianion a t high pH, the solution was buffered between p H 4 and 6:
- H+
Hi
I
1OH210P03H2
I
C10H210P03H-
Cl,H,10P0,H2-
i H+
The first and second pK,'s of butyl phosphate are 1.8 and 6.84, re~pectively,~' and we assume that the pK,'s of decyl phosphate (24) Romsted, L. S.; Zanette, D. J . Phys. Chem. 1988, 92, 4690. (25) Funasaki, N . J . Phys. Chem. 1979, 83, 237. (26) Hautecloque. S.; Grand, D.; Bemas, A . J . Phys. Chem. 1985, 89, 2705. (27) Bunton, C. A,; Cerichelli, G . Inr. J . Chem. Kinet. 1980, 12, 519. (28) Dunlap, R . B.; Romsted, L. R.; Cordes, E. H. J . Phys. Chem. 1967. 71. 458 I . (29) Bunton, C. A.; Robinson. L. J . Phys. Chem. 1970, 74, 1062. (30) Dorion, F.; Charbit. G.: Gaboriaud. R. 3. Colloid Interface Sci. 1984, 101. 2 7 . (31) Kumler, W . D.;Eiler. J . J . J . Am. Chem. SOC.1943, 45, 2355.
Experimental Section Materials. The preparation and purification of PADA, NaLS, decyl phosphate, and N a D P were described p r e v i ~ u s l y .The ~~ alkali-metal monoanion salts of M D P (KDP, RbDP, and CsDP) were prepared by titrating warm, ethanolic monodecyl phosphate solutions with 1 equiv of aqueous, standardized solutions of approximately t M KOH, RbOH, and CsOH, respectively. The products, which crystallized on cooling overnight in the refrigerator, were isolated and dried under vacuum. Lithium decyl phosphate, LiDP, was prepared by the same method, but it proved to be insoluble, even in almost boiling water, making it unsuitable for micellar studies. All inorganic salts and succinic acid were reagent grade and, except for vacuum drying, were used without further purification. All solutions were prepared from doubly distilled, demineralized, carbon-filtered water that was boiled under nitrogen to remove dissolved carbon dioxide. Methods. Absorbance measurements were carried out on a Perkin-Elmer 559A UV/vis dual-beam spectrophotometer a t 50 f 0.1 OC. W e used 0.08 M M D P and 0.04 M N a L S in all indicator experiments to ensure that PADA was completely micellar bound.24 The same solution preparation protocol was used in each salt effect study. Small aliquots (10-50 pL) of a concentrated solution of MC1 (4-5 M) were added successively to a 2.2-mL solution containing surfactant, succinate buffer prepared to give a preset initial pH, pH,, and 4-6 pL of 0.01 M (32) Okamoto. Y. Bull. Chem. SOC.Jpn. 1985, 58, 3393 and references therein. (33) Arakawa, J.; Pethica, B . A. J . Colloid Interface Sci. 1980, 75, 441. (34) Tahara. T.; Satake, I.; Maturra, R. Bull. Chem. SOC. Jpn. 1969, 42, 1201. (35) Chevalier, Y.; Chachaty, C. Colloid Polymer Sci. 1984, 262, 489. Kamo, 0.;Matsushita, K.; Terada, Y.; Yoshida. T.; Okabayashi, H.; Chem. Scripta 1984, 23, 189. Chevalier, Y.; Belloni, L.; Hayter, J. B.; Zemb, T. J . Phys. (Paris)1985, 44, 749. Chevalier, Y.; Chachaty, C. 3. Phys. Chem. 1985. 89, 875. Chevalier, Y . ;Chachaty, C. J . Am. Chem. SOC.1985, 107, 1102. Chachaty, C.; Ahlnas, T.; Lindstrom, B.; Nery. S . H.; Tistchenko. A . 31. J . Colloid Interface Sci. 1988, 122. 406.
The Journal of Physical Chemistry, Vol. 93, No. 10, 1989 4221
indicator Equilibria in Micellar Solutions solution of PADA in CH3CN (spectrophotometric grade). Because the absorbance of PADA decreases with increasing pH, the final PADA concentration is increased at higher pHi to maintain the precision of the absorbance readings. The net absorbance, i.e., sample solutions versus reference of identical composition minus PADA, was recorded after each addition of MCI. The values of pH, were about 4.4, 4.8, 5.2, 5.5, and 5.9. The highest concentration of added MCI was 0.4 M. The indicator ratio of neutral, A, to protonated, AH, forms of PADA was calculated by using eq 1 from the measured absorbance values at X = 552 nm in NaLS and X = 545 nm in MDP, after correction for dilution (greatest correction about 8%): [AI / [AH1 =
(Amax
- A ) / ( A - Amin)
Theory Equation 2 describes the effect of anionic micelles on the in[AH1
@H
(2)
dicator ratio of completely micellar bound PADA. The complete derivation, following the assumptions of the PIE model, was published earlier and is only summarized here.24 At constant total surfactant concentration, [DT], the concentration ratio of the two forms of the indicator is directly proportional to the ratio of counterion activity, aM,to hydrogen ion activity, aH,in the aqueous phase, where is the degree of counterion binding, KHMis the ion-exchange constant between the alkali-metal counterion and the hydronium ion, and KAmis the intrinsic acidity constant of PADA in the micellar pseudophase. Equation 2 is obtained by combining eq 3-5, where mHS= KA"' = [A]mHs/[AH] K ~ M=
aHmMS -
[Hwl [MmlYHmYMm
aMmHS
[Mwl [HinhMwYH" mMS = p
I
(3)
(4) (5)
UH"/([DT] - cmc) and mMS= aMs/([DT]- cmc). The activities of the micellar-bound proton, mHm,and the alkali-metal counterion, mMm,are expressed as unitless ratios of micellar-bound
1
d
6
2 - 0
4
\ 2
L
(1)
where A,,, and A,,, are the absorbances of the fully protonated and deprotonated species, respectively. Square brackets indicate concentration in moles/liter here and throughout the text. We measured the molar absorptivity, e , of the acid form of PADA in KDP, RbDP, and CsDP and found that e is independent of counterion type and numerically the same as in N a D P within We assumed that e is also independent of counterion type in NaLS. In a separate set of experiments, using the same protocol for addition of MCI, we recorded the solution pH at 50 OC a t every MCI concentration used in the indicator experiments. However, the initial solution volume was 5 mL, instead of 2.2 mL, and the aliquot size was adjusted so that, after each addition, the final concentrations of all components were identical in the indicator and pH (except [PADA] = 0) experiments. We used a Ross semimicro combination pH electrode, Model 81-03, in a stirred vessel, thermostated at 50 f 0.1 OC, and a Corning pH meter, Model 130. The electrode was calibrated against standard pH 4.0 and 7.0 buffers before each run. Reproducibility was very high, generally 1 0 . 0 1 of a pH unit. Critical micelle concentrations, cmc's, were also measured at 50 f 0.1 "C by using a Fisher DuNouy tensiometer a t each pH, for each counterion and at four to five salt concentrations at pH, = 5.2 (Table I). Our cmc values for NaDP and KDP in 0.02 M succinate buffer are in good agreement with literature values obtained in the absence of added salt.33 log cmc versus log MCI plots were linear. iterative least-squares fits of the data24were used to calculate the cmc a t each concentration of added salt. Note that the cmc for MDP is not very sensitive to counterion type.
-[AI - -KA~KH~~M
- 4
0
0
2
6
10
14
-4
(ocs/oH)xlo
Figure 2. Effect of added CsCl on the indicator ratio of PADA in 0.08 M CsDP, 0.02 M succinate buffer at 50 OC, plotted against three different activity ratios of Cs' and H+in the aqueous phase. Values of the
indicator ratios are the same for all three curves, and the slopes are similar, so the abscissas are offset for clarity. The activity ratios are as follows: acS/aH(C~), curve A; Ucs/UH(Na),curve B; acs/aH(Na),which
includes NaDP concentration correction (see text), curve C. Solid lines are least-squares fit of the data. Each symbol represents a set of CsCl additions at a different pHi (see text); (m) 4.402; ( X ) 4.820; (0) 5.200; (0)5.500; (0) 5.918.
=i
/ RbCl
51
/ 1
\
a
u
1
0 0
2
4
6
Figure 3. Plots of indicator ratios of PADA versus metal ion and hydronium ion activity ratios in the aqueous phase in solutions of 0.08 M
MDP, 0.02 M succinate buffer, for the four alkali-metal salts. Symbols have the same meaning as in Figure 2. Solid lines are calculated by a least-squares fit; the line for CsCl is from Figure 2 and the line for NaCl from Figure 3 in ref 24. The abscissa for RbCl and CsCl is offset for clarity. ion to micellized surfactant ( [DT] - cmc) to avoid having to define their location, or reactive volume element, within the micellar pseudophase.6*21.22 The monomer concentration in the aqueous pseudophase is assumed to be equal to the cmc. We note that KAm is also a unitless number because of the definition of the concentration of bound protons, but it can be converted to more conventional units by defining a volume element for the location of the indicator within the micellar pseudophase (see Discussion). Because the alkali-metal ion concentration is always 2-3 orders of magnitude greater than the proton concentration and KHMis on the order of 1 (see Discussion), we assume that the fraction of the micelle surface covered by associated but not covalently bound protons is negligible (i.e., mHS N a > K > R b > Cs, which correlates with the size of the naked cation, Le., the smaller the cation the stronger the interaction, is associated with a high field strength generated by the surface charges partially or completely dominating the hydration of the cations. Whereas Eisenman’s series I, Cs > Rb > K > Na > Li, which correlates with the size of the hydrated cation, Le., the smaller the hydrated cation the weaker the interaction, is associated with a low surface field strength too weak, or to dispersed, to overcome the hydration of the cations. (48) Strauss, U. P.; Leung, Y. P. J . A m . Chem. SOC.1965, 87, 1476. Armstrong, R. W.; Strauss, U. P. In The Encyclopedia of Polymer Science and Technology; Bikales, N. M., Ed.; Wiley-Interscience: New York, 1969; Vol. 10, p 781. (49) Eisenman, G. In The Glass Electrode; an Interscience Reprint, Interscience: New York, 1965; p 213. (50) Diamond, J. M.; Wright, E. M. Ann. Reu. Physiol. 1969, 31, 581. ( 5 1 ) Diamond, J . M. J . Exp. Zool. 1975, 194, 227.
The Journal of Physical Chemistry, Vol. 93, No. 10, 1989 4225 A number of properties of both cationic and anionic micelles are correlated with the size of the hydrated counterion, Le., they follow a Hofmeister series,44including cmc’s, aggregation numbers and degrees of counterion binding,41,52and ion-exchange constants, reaction rates, and indicator equilibria.6~21~22 Of our five M D P surfactants, only LiDP is completely insoluble, even in boiling water, indicating that the solid state is very stable and the total dominance of electrostatic over hydration forces. However, with all of the other four counterions, once the surfactant is in solution, the selectivity order depends on the size of the hydrated cation. To date, no micellar property is known to correlate with the ionic radius of the counterion, indicating that hydration of head groups and counterions is probably a fundamental requirement for micelle stability. Ionic surfactants spontaneously form thermodynamically stable solutions of dynamic aggregates, unlike vesicles which are metastable aggregates that fuse and precipitate slowly. Indeed, the term “pseudophase”, commonly used as a conceptual picture for interpreting many of the physicochemical properties of micellar solutions, reminds us of the micelles are not a true separate phase. Clear examples of site binding with micelle-forming surfactants are associated with precipitation. For example, in moderately acidic solutions (e.g., pH Il), the premier site-binding cation, the proton precipitates micellar solutions of surfactants with weakly acidic head groups such as long-chain alkyl carboxylates and phosphate^,^^ because the proton binds covalently to the head group and neutralizes its charge. But surfactants with strongly acidic head groups that have protons as their only counterion, such as alkyl sulfates,53 are very soluble. The alkyl sulfonates remain in solution even in the presence of excess acid.s4 The near unity of the ion-exchange constant between sodium ions and the proton in NaLS45is consistent with the proton retaining its water of hydration and being associated, but not covalently bound a t the micelle surface. Another intriguing possibility is that the reversal in selectivity order is related to the surface curvature of the aggregate.s5 The high curvature of the micelle surface reduces the surface charge density, inducing selectivity order I, compared to the large, nearly flat, more densely packed surfaces of vesicles which generate a higher field strength and order XI. Mortara et al. found that the flocculation rates at 50 OC of sonicated vesicles of dihexadecyl phosphate depended upon the type of alkali-metal salt added and the rates decreased in the order Cs > Na > K > LLs6 Although this series is not part of any of Eisenman’s 11 orders for the five alkali-metal cations, it is almost the same as our order for MDP. And unlike the probably large vesicles used in the work listed in Table IV (vesicle size was not specified), the dihexadecyl phosphate vesicles were known to be small single-walled vesicles (diameter about 500 A) with relatively high radius of curvature. Conclusions
Three main conclusions can be drawn from this work. First, the PIE model provides a complete description of alkali-metal counterion effects on the indicator ratio of PADA bound to MDP micelles. Second, when the intrinsic acidity constant of micellar bound PADA is expressed in ordinary units of moles/liter of Stern layer volume, its value is close to that in water, consistent with the view that the micelle surface is very wet. Third, assuming /3 and KAmare independent of counterion type, the selectivity order for counterion binding is the same for M D P and MLS micelles: Cs > R b > K > Na. To our knowledge, this is the first estimate of the selectivity order of MDP toward monovalent alkali-metal counterions and it is the reverse of the selectivity order shown by (52) Kresheck, G. C. In Water, A Comprehensive Treatise; Franks, F., Ed.; Plenum Press: New York, 1975; Vol. 4, p 95. (53) Gaboriaud, R.; Charbit, G.;Dorion, F. J . Colloid Interface Sci. 1984, 98, 583. (54) Bunton, C. A,; Romsted, L. S.; Savelli, G. J. Am. Chem. SOC.1979, 101, 1253. ( 5 5 ) Evans, D. F. Langmuir 1988, 4, 3. ( 5 6 ) Mortara, R. A,; Quina, F. H.; Chaimovich, H . Biochem. Biophys. Res. Commun. 1978, 81, 1080.
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The Journal of Physical Chemistry, Vol. 93, No. IO, I989
TABLE V: Parameters Used To Estimate in Monoanion Form CSCI, acid pKaQ pKAV6 pKAmc M ROP0,H- 6.84 6.34 7.17 0.402 1.3 2.15 0.0 ROP03H, 1.8
the Fraction, f, of CsDP
pH 5.681 3.89 4.878 2.21
mHS
X X
f 0.98 0.97
First and second pKa’s of butyl phosphate.29 ‘pKA’ = pKa - 0.5. Equation 3 applied to acidity constants of decyl phosphate. dUsing eq 3, [ A H ] / [ A ] = 0.1 12. eUsing eq 3, [ A H ] / [ A ] = 0.637.
several different anionic phospholipid vesicles.
Acknowledgment. We thank the Research Council and Biological Sciences Research Fund of Rutgers University, the National Institutes of Health, GM32972, the National Science Foundation U.S.-Latin American Cooperative Program, Brasil, the donors of the Petroleum Research Fund, administered by the American Chemical Society, and the Research Corp. for financial support. We also thank Hernan Chaimovich and Frank Quina, Universidade de Sao Paulo, Faruk Nome, Universidade de Santa Catarina, Clifford Bunton, University of California, Santa Barbard, and U. P. Strauss of this department for very helpful discussions and Olimpiu Petrescu for rescuing the file containing an almost complete version of the manuscript from computer oblivion. Appendix I Assuming all the error was in the absorbance readings and the indicator ratio a t [A]/[AH] = 2.75 (see Figure 2, curve A, pHi = 5.918, first circle) should actually fall on one of the lower curves, e.g., squares of diamonds, lines, then the ratio should be 2.0 (squares) or 1.5 (diamonds). This is a substantial decrease and would require the measured absorbance to be 2 0 4 0 % larger, much greater than our experimental error. A second possibility is that [M,] is too low (eq 7). The major sources of uncertainty in [M,] are the cmc and p. To make the same datum, aCs/aH= 2.9, coincide with the least-squares line, the ratio would have to be 4.5, a 72% increase in acs, This would require that the cmc be too low by a factor of 2.9, a highly unlikely possibility, or an increase in ( 1 - /3) from 0.3 to 4, i.e., the micelle would be 400% ionized, a nonsensical result. For the following reasons we believe the dispersion in our data for added CsCI. RbCI, and KCI is caused by the sensitivity of the glass electrode to these ions and that the most accurate pH readings are in the presence of added NaCI. (a) In Figure 5, the pH-log [ MCI] curves for NaCl are consistently lower than the curves for KCI, RbCI, and CsCl in M D P and succinate buffer (Figure 5 A ) , succinate buffer (Figure 5B), and dilute HCI (Figure 5C57). The striking similarity of these patterns shows that the pH difference between NaCl and the other three salts cannot be caused by surfactant or buffer effects. Exactly the Same patterns were obtained at pH, = 4.4, 4.8, 5.5, and 5.9 in 0.08 M MDP and 0.02 M succinate buffer (results not shown). (b) Plots of [A]/ [AH] versus aNa/aHfor added NaCl show the least dispersion (Figure 3 in ref 24). (c) The measured pH values in 0.04 M NaLS. 0.02 M sodium succinate buffer, pH, = 5.705 (not shown), show the same pattern as in MDP, indicating that the pattern does not depend on surfactant type. (d) The measured indicator ratios of PADA in 0.02 M succinate buffer, p H i = 5.2, up to 0.4 M MCI for all four salts (Figure 5D) are essentially the same (maximum average deviation &6%) with no difference between NaCl and (57) The results in Figure 2C were obtained by adding small aliquots of 4 or 5 M MCI solutions containing 0.001 M HCI to a stock solution of 0.001 M HCI.
He et al. other three salts. Indeed, the average total change in the indicator ratio, h log [AH]/[A], is 0.22 f 0.02 for all four cations. This value is close to the total pH change for added NaCI, ApH = 0.24, but significantly different from the average pH change for addition of the other three cations, h p H = 0.13 f 0.013. This result strongly supports our conclusion that the specific salt effects are on the electrode and not the buffer or, in this case, the indicator. (e) Finally, a Brinkmann glass electrode with a Ag/AgCl reference electrode gave the same pattern of pH response, although slightly different pH values (not shown), as the Ross electrode under identical conditions, 0.02 M succinate buffer, pHi = 5.2 (Figure
5B).58 In sum, the pH-log [MCI] profiles in the presence of buffer (e.g., Figure 5B) and buffer plus surfactant (e.g., Figure 5A) must be a composite of several effects. Added salt reduces the activity coefficient of the hydronium ion (with a modest dependence on alkali-metal ion type),37and together with the specific salt effect on the glass electrode (Figure 5C) would, by themselves, increase solution pH. However, this p H increase must be overwhelmed by the general salt effect on the ionization of the buffer, probably on both its first and second pK,’s, because added salts always reduces the pH in the presence of succinate buffer. Adding MDP also increases the measured pH,24 because both micellar and monomer decyl phosphate monoanion are weak acids, but the salt effect pattern in not affected (compare curves A and B in Figure 5). Appendix I1 The fraction of CsDP in monoanion form a t the extremes of pH and [CsCI] (Le., the two terminal data points in Figure 2, curve C ) can be calculated from the indicator ratio by using eq 3, if the surface acidity, mHS,and the intrinsic PKA”s for the first and second ionization of decyl phosphate micelles can be estimated. The intrinsic pKAm’swere obtained by assuming that micellization affects the pK, for the first and second ionization of micellized CsDP to the same extent as the pK, of PADA (about 0.5 pK, units lower than the value in water). Thus pK,’ = pK, - 0.5, where pK, is the acidity constant for the first and second ionization of butyl phosphate in water. The PKA”’s are then converted to pKA”s by using eq 1 1 and the same value for V,. Values for mHSwere calculated by using eq 3, KA”’ for PADA in M D P (Table I l l ) , and the measured indicator ratios for the two extreme data points. The results, summarized in Table V, indicate that CsDP is more than 95% in its monoanion form a t the extremes of our data, supporting our assumption that MDP micelles are always in their monoanion form. Complete confirmation requires an independent estimate of the intrinsic pK,’s of MDP, which we hope to obtain by 3 1 PN M R or FTIR. The interfacial pH, Le., the concentration of “associated” but free protons, was estimated by using the relation aHm= mHm([DT] - cmc) at the same two set of conditions for CsDP listed in Table V. At solution pH = 4.878, [CsCI] = 0.0, [CsDP] = 0.08 M, the Stern layer pH = 2.77, a 130-fold increase in local proton concentration. However, at solution pH = 5.681, [CsCI] = 0.4 M, [CsDP] = 0.08 M, the Stern layer pH = 4.52, only a 14.5-fold increase in local proton concentration. Like apparent acidity constants, interfacial pH in micellar solutions depends on a number of factors including surfactant concentration and type, solution pH, and counterion concentration and type. Registry No. PADA, 13103-75-8;SDS, 151-21-3; KDP, 68427-32-7; RbDP, 119696-16-1; CsDP, 119696-17-2; NaDP, 60160-24-9: KaCI. 7647-14-5; KCI, 7447-40-7; RbCI, 7791-1 1-9; CSCI, 7647-17-8. (58) For example, at 0.0249 Madded NaCI. the pH reading is 5.219 with the Ross electrode and 5.169 with the Brinkmann electrodes.
