J. Phys. Chem. 1992, 96, 1463-1467 is facilitated by increasing quencher concentration. The decrease of the second-order rates with increasing global concentration [Q], is a particularity of the restricted reaction space.6 An increase in [Q] means a decrease in the available free sites per quencher. Thus, even though the quenching species becomes more abundant, facilitatingfmt-order kinetics, the probability for a single quencher to reach the fluorophore decreases, making second-order kinetics slower. The k's do not always decrease with increasing [Q], Thus for pyrene in mPG vesicles, which provide a less restrictive lipidic core, it has been found that both K's and k's increase with [Q]. However, in that case, the increase in the K's is faster. Notice also, that for DPH in the DPPGLcYPGmixture in Table I1 and pyrene in DPPG in Table I11 kl increases with [Q]. We attribute this last result to the high efficiency of the reaction at short times and high quencher concentrations, due to quasi-static quenching. Failure to repeat this same result with the pair pyrene-12-DSME in DPFG in Table I11is most probably due to the limited accuracy of the data obtained with this last system. Notice that both first-order and second-order rates are 1 order of magnitude higher for all data obtained with DPH than for those obtained with pyrene. This difference in reaction efficiency goes along with the time scales of survival of the two excited species. In other words, the reaction is faster for the fast-decaying DPH. Of course, this behavior is expected when a substantial quenching occurs within the same concentration ranges. Notice the large difference between the reaction rates at short times (KI) and at long times (KJ. This is due to the relatively low f values observed in this work (cf. Figure 1). The average R is, generally, higher in L ~ P G than in DPPG vesicles. Apparently, this is due to the higher fluidity of the LCYPGlipidic core. Notice, finally, the extensive increase of K1with [Q] in the vesicles of the DPPG-mPG mixture (Table I) and of DPPG (Table 11). This is consistent with our above assertion of quasi-static quenching, in particular in these systems. The above discussion shows that the information on reaction rates obtained through eqs 1 and 2 provides an extensive knowledge of the behavior of the reactants in microheterogeneous environments.
1463
An additional final remark concerns the behavior under variable temperatures. Previous measurements with pyrene excimers have shown that an increase in temperature increases the calculated fvalues? However, no variation infwas detected with DPH. We attribute this result to the time scale of the DPH data, which is too short to allow any detectable temperature effect.
Conclusion The fluorescence decay of DPH embedded in small unilamellar vesicles of DPPG, L ~ P Gor, a mixture of these two phospholipids (80% DPPG, 20% LCXPG) has been studied in the presence of varying concentration of 12-DSME used as fluorescence quencher. The vesicle concentration was held constant. The results have been compared with similar results obtained with pyrene in the present and previous works. All decay profiles have been successfully fitted with eq 1 through its form in eq 3. A decay rate has been defined through eq 2. The results of the analysis are easy to explain, and they help in redefining the reaction rate in organized lipid vesicles by introducing the concept of dimensionality through the exponent$ The analysis presented permits drawing conclusions on the probe distribution in the phospholipid bilayer of the vesicles and the possibility of fast kinetics between the reacting species (i.e. quasi-static quenching). The geometry of the vesicle lipidic core is not "seenn the same way by different fluorophorequencher pairs, i.e. the f values depend on the considered pair for the same vesicles. Mixtures of DPPG and mPG lipids give vesicles with rather particular structures which are detected only by the present method. 12-DSME quenchers DPH fluorescence much faster than it quenchers pyrene fluorescence. Acknowledgment. We acknowledge help from the Technical Staff of LURE and Dr. F. Merola, who is in charge of the SA1 beam and the experimental set up. We also acknowledge financial aid from an EEC-LURE grant. Registry NO. DPH, 1720-32-7; 12-DSME, 137917-39-6; DPPG, 4537-77-3; LaPG, 6418-92-4; pyrene, 129-00-0.
Effects of Monoalkyl Phosphate Surfactants upon the Acid Hydrolysis of Dioxolanes Angelo A. Ruzza, Maria Rosfinia K.Walter, Faruk Nome, and Din0 Zanette* Departamento de Quimica, Universidade Federal de Santa Catarina, 88049 Florian@olis, SC, Brazil (Received: May 29, 1991; In Final Form: August 28, 1991)
The pseudophase ion-exchange (PPIE) model was tested in the presence and absence of buffer for the acid-catalyzed hydrolysis of 2-(pmethoxyphenyl)- 1J-dioxolane @-MPD)and 2-(2,4-dimethoxyphenyl)-2-ethyl-1,3-dioxolane(2,4-DPED) in the presence of sodium monoalkyl phosphate surfactants. The maximum increase in the observed rate constant was about 10-fold for both substrates. The second-order rate constant for the acid hydrolysis of 2,4-DPED in the micellar phase of the phosphate surfactants (k2m= 14.3 M-'s-l) is similar to that observed in the bulk aqueous phase (& = 12.7 M-I s-l). WhenpMPD was used as a substrate, a modest 3-fold enhancement in the second-order rate constant is found. The application of the PPIE model requires careful pH measurements, since the surfactant itself is a weak acid.
Introduction The pseudophase ion-exchange (PPIE) formalism has been used to interpretate the effects of charged interphases on reaction rates of a number of systems such as aqueous ionic micelles,'-' synthetic vesicles,5-8reversed micelles? and microemulsions.'O.ll Anionic
and cationic surfactants have a significant effect upon indicator equilibria1"l9 and bimolecular reaction^,^-^**,*^^ including specific W. L. J. Am. Chem. Soc. 1981, 103, 5439. ( 6 ) Cuccovia, I. M.; Quina, F. H.; Chaimovich, H. Tetrahedron 1982, 38, ( 5 ) Fendler, J. H.; Hinze, 917.
(1) Fendler, J. H.; Fendler, E. J. Catalysis in Micellar and Macromolecular Systems; Academic Press: New York, 1975. (2) Quina, F. H.; Chaimovich, H. J . Phys. Chem. 1979.83, 1844. (3) Romsted, L.S. Ph.D. Thesis, Indiana University, 1975. (4) Berezin, I. V.;Martinek, K.;Yatsimirski, A. K. Russ. Chem. Rea (Engl. Trawl.) 1973, 42, 487.
(7) Chaimovich, H.; Bonilha, J. B. S.,; Zanette, D.; Cuccovia, I. M.In Surfactants in Solution; Mittal, K. L.,Lindman, B., Eds.; Plenum: New York, 1984, p 1121. (8) Fendler, J. H. Membrane Mimetic Chemistry; Wiley: New York, 1982. (9) El Seoud, 0.A.; Chinelatto, A. M. J . Colloid Interface Sci. 1983, 95,
163.
0022-365419212096-1463$03.00/0 0 1992 American Chemical Society
1464 The Journal of Physical Chemistry, Vol. 96, No. 3, 1992
salt Usually, the PPIE formalism provide8 qualitative and quantitative interpretations of these effects without complex equations or computational efforts. The validity of the PPIE formalism has been tested repeatedly, and the model fails in extreme conditions, mainly with hydrophilic reactive counterion surfactants22and with normal micelles in the presence of high concentrationsof H+ or OH-.23-27 Recently, it has been demonstrated that the apparent failure for hydrophilic reactive counterion surfactants can be solved by considering changes in the degree of dissociation as a function of detergent concentratione2' Alkyl phosphate surfactants are used in many industrial applications, such as laundry, emulsification, and mineral flotation, besides being used as antistatic agents and corrosion inhibitors. Micelles with a phosphate head group have been suggested to mimic the complex interphase of biological membranes, since they are more stable and have less complex structures.8 Although descriptions of synthetic preparations and of physical properties of some alkyl phosphate m i ~ e l l e s ~ have ~ - ~ lbeen published, little is known about the chemical and interfacial properties of these surfactants. Recently, Romsted and Zanette15 published studies about the acidity of the micellar surface of sodium decyl phosphate monoanion, NaDP, evaluated using the indicator pyridine-2-azo-p dimethylaniline, PADA, as an interfacial probe, at several pH values and surfactant concentrationsover a large NaCl concentration range. They reported the effect of added univalent counterions (Na, K, Rb, Cs)up to 0.4 M on the indicator ratio of PADA in micellar solutions of alkali-metal (Na, K, Rb, Cs) salts of decyl phosphate.16 Some important conclusions were as follows: (a) the PPIE model provided a complete description of alkali-metal counterion effects; (b) the selectivity order for counterion binding is the same for the alkali-metal decyl phosphates and dodecyl sulfate micelles, that is, Cs > Rb > K > Na; (c) no special effects of the head groups of micellized decyl phosphate were found in the indicator equilibria when compared with other well-known anionic surfactants like sodium dodecyl sulfate (SDS). In this paper, we report the catalytic effects of sodium decyl and dodecyl phosphate on the acid-catalyzed hydrolysis of 2-(p(10) Mackay, R. A. J . Phys. Chem. 1982,86,4756. (11) (a) Ricardo, D. R. P.; Zanette, D.; Nome, F. J. Phys. Chem. 1990, 94,356. (b) Silva, I. S.;Zanette, D.; Nome, F. Arual. Fis.-him. Org., Conf., 4th 1985, 123. (12) Chaimovich, H.; Aleixo, R. M. V.; Cuccovia, I. M.; Zanette, D.; Quina, F. M. In Solurion Behavior of Surfactants: Theoretical and Applied Aspects; Mittal, K. L., Fendler, F. J., Eds.; Plenum Press: New York, 1982; Vol. 2, p 949. (13) Zanette, D.; Leite, M. R.; Reed,W.; Nome, F. J . Phys. Chem. 1987, 91, 2100. (14) Fernandez, M. S.; Fromherz, P. J . Phys. Chem. 1977, 81, 1755. (15) Romsted, L. S.; Zanette, D. J. Phys. Chem. 1988,92,4690. (16) He, Z. M.; OConnor, P. J.; Romsted, L. S.;Zanette, D. J . Phys. Chem. 1989,93,4219. (17) Drummond, C. J.; Grieser, F.; Healy, T. W. J . Phys. Chem. 1988, 92, 2604. (18) Romsted, L. S. J. Phys. Chem. 1985,89, 5107. (19) Romsted, L. S. J . Phys. Chem. 1985,89, 5113. (20) Bunton, C. A.; Savelli, G. Adu. Phys. Org. Chem. 1986, 22, 213. (21) Bunton, C. A.; In Reaction Kinetics in Micelles; Cordes, E . H., Ed.; Plenum Press: New York, 1973; p 73. (22) Gonsalves, M.; Probst, S.; Rezende, M. C.; Nome, F.; Zucco, C.; Zanette, D. J . Phys. Chem. 1985, 89, 1127. (23) Bunton, C. A.; Romsted, L. S. In Solurion Behavior of Surfacrcmrs: Theorerical and Applied Aspecrs; Mittal, K. L., Fendler, E. J., Eds.;Plenum: New York, 1982; p 975. (24) Nome, F.; Rubira, A. F.; Franco, C.; Ionescu, L. G. J . Phys. Chem. 1982, 86, 188 1. (25) Romsted, L. S. In Surfactants in Solurions; Mittal, K. L., Lindman, B., Eds.; Plenum: New York, 1984; Vol. 2, p 1015. (26) Stadler, E.; Zanette, D.; Rezende, M. C.; Nome, F. J . Phys. Chem. 1984,88, 1892. (27) Nwes, M. F. S.; Zanette, D.; Quina, F.; Moretti, M. T.; Nome, F. J . Phys. Chem. 1989, 93, 1502. (28) Nelson, A. K.; Toy, A. D. F. Inorg. Chem. 1963, 2,775. (29) Imokawaa, G.; Tsutsumi, H. J. Am. 011 Chem. Soc. 1978,55,839. (30)Tahare. T.;Satake, I.; Matuura, R. Bull. Chem. Soc. Jpn. 1%9,42, 1201.. (31) Arakawa, J.; Pethica, B. A. J . Colloid Interface Sci. 1980, 75, 441.
Ruzza et al. TABLE I: First-Order Rate Colrptiatr .(Ia F d o a of tbe Hydrogen Ion Coarmtmtiolrp for tbe Hydrolyses in Water of p-MPD .ad 2,4DPED at 50 O C in 0.01 M Formate Buffer 104[H+], 1 0 3 k b , w 9 s-' M" p-MDP 2,4-DPED 6.30 121.0 8.16 3.89 78.0 4.46 2.50 51.0 3.00 1.58 31.0 2.19
1@[H+], 103kb,w,s-' Ma p-MDP 2.4-DPED 1.00 18.0 1.38 0.63 12.0 1.11 0.40 1.0 0.64
Hydrogen ion concentrations were calculated using the measured pH values, ignoring activity coefficient changes.
methoxypheny1)-1,3-dioxolane (p-MPD) and 2-(2,4-dimethoxyphenyl)-2-ethyl-1,3-dioxolane (2.4-DPED) in the presence and absence of succinate buffer. Complete treatment of the kinetic data using the PPIE formalism was only possible by measuring the hydronium ion concentration in the aqueous phase. The acidity of the NaDP aqueous solutions changes as a function of surfactant concentration, probably because the surfactant itself is a weak acid. This behavior of the monoalkyl phosphates is strikingly different from other anionic surfactants such as SDS.
Experimeotal Section Materials. Decyl phosphoric (DPA) and dodecyl phosphoric (DDPA) acids were prepared from ldecanol (Aldrich, 99%) and 1-dodecanol (Merck), respecti~e1y.I~The products were recrystallized extensively from hexane giving white crystalline solids: mp 48 OC (lit. 49 OC31and 45 OCB) and 58 OC (lit. 58 0C28)for DPA and DDPA, respectively. Sodium monodecyl phosphate (NaDP) and sodium monododecyl phosphate (NaDDP) were prepared by neutralization of a methanolic solution of the corresponding acid with standardized 1 M NaOH. The products were precipitated by adding excess of alcohol to the final methanol/ water solution and then isolated and dried under vacuum. The critical micellar concentrations (cmc) for NaDP and NaDDP were obtained from plots of specific conductivity versus surfactant concentrations at 50 OC giving values of 0.036 M for NaDP (literature 0.0364 M at 60 "C3') and 0.0114 M for NaDDP (literature 0.0094 M at 55 0C32). 2-(p-methoxyphenyl)-1,3dioxolane (p-MPD) was prepared according to the procedure of Fife and JaoUJ32-(2,4-dimethoxyphenyl)-2-ethyl-1,3-dioxolane (2,4-DPED) was prepared according to the procedure of D. R. White." The purities of both substrates (pMPD and 2,4-DPED) were found to be satisfactory by TLC, UV,and IR analyses. All solutions were prepared using distilled, demineralized water, which was boiled and cooled under nitrogen to remove carbon dioxide and kept under a nitrogen atmosphere. All other reagents were the best available reagent grade and were used without further purification. Methods. The formations of the products p-methoxybenzaldehyde and 2,Cdimethoxyphenylpropiophenone were followed spectrophotometrically at 284 and 266 nm, respectively, using a Shimadzu Model UV 210-A spectrophotometer equipped with a thermostated water-jacketed cell compartment. The kinetic data were directly stored in a microcomputer, through a Microquimica 12 bits A/D interphase board, and rate constants were calculated using an iterative least-squares program. Plots of In (A, - A,) vs time were linear for at least 90% of the plots, with correlation coefficients greater than 0.99. All pH measurements were carried out using a MICRONAL pH meter (Model B-222) which was calibrated prior to use with standard buffers of pH = 4.00 and 7.00. The experimental pH of all kinetic solutions was measured, since the surfactant con(32) The value of cmc, taken from ref 3 1, amtsponds to potassium dodecyl phcsphate at 55 "C. The same reference reports cmc valuts of 0.0355 and 0.0364 for potassium and sodium decyl phosphate at 60 'C, respectively, indicating that the counterions have a negligible effect on the cmc. (33) Fife, T. H.; Jao, L. K. J. Org. Chem. 1965,30, 1492. (34) Glatz, B.; Helmchen. G.; Muxfeldt, H.; Pocher, H.; Prewo, R.; S ~ M , J.; Stezowski, J. J.; Stojda, R. J.; White, D. R. J. Am. Chem. Soc. 1979. 101, 2171.
Surfactant Effects on the Acid Hydrolysis of Dioxolanes
0
5.0
10.0
The Journal of Physical Chemistry, Vol. 96, No. 3, 1992 1465
0
15.0
10.0
5.0
lO*[NaDP], M
Figrw 1. Plot of observed rate constants for the acid-catalyzed hydrolysis of p-MPD in aqueous solutions of NaDP at 50 OC.
centration affected solution acidity. The alkyl phosphate surfactants are weak acids, their dissociation equilibrium being affected by both ionic strength and specific i o n ~ . ' ~ J ~ Kinetic measurements for the determination of the effect of NaDP and NaDDP on the acid hydrolysis of p M P D ad 2,4DPED were started by addition of 10 pl of a stock solution of the M in acetonitrile) into 2.5 mL of the appropriate substrate ( aqueous surfactant solution.
F i i 2. Plot of observed rate constants for the acid-catalyzed hydrolysis of 2,4-DPED in aqueous solutions of NaDP at 50 OC.
7
n o'C'_oRl
C-R1
R3
20
u)
Results Table I shows first-order rate constants in water (kob,w)at different [H+] for the hydrolysis reaction of the substratesp M P D and 2,4-DPED (eq 1). Under these experimental conditions, both 0
0 0
5
II
reactions are first order in organic substrate and in hydrogen ion concentration. The overall second-order rate constants (kZJ were calculated from the slope of the linear plots (not shown) of versus [H+], giving values of 12.2 and 194 M-' s-' for 2,4-DPED and p-MPD, respectively. Measurements of the pH of the aqueous solutions containing the phosphate surfactants show an increase in the solution pH as a function of increasing surfactant concentration (Table 11). This behavior, a rise of the pH values beginning at the cmc, has been observed with alkyl phosphates and can be almost certainly attributed to exchange of hydrogen ions from the aqueous phase with sodium ions in the micelles. Attempts to control the solution pH by adding 20 mM sodium succinate/succinic acid buffer were not completely successful although the buffer did reduce the extent of the pH change (compare columns 4 and 5, Table 11). In this pH range, succinic acid is present only in its mono and dianion forms, which should preferentially be in the aqueous phase and be excluded from the interfaces of the anionic micelles. The ApH = 0.4 (column 5, 5.68 - 5.28) shows that a substantial fraction of the total hydrogen ions (free, bound to buffer, and bound to micelles) is being incorporated into the micelles. Figures 1 and 2 show the effect of the NaDP surfactant upon the observed rate constants (kobs,J of the acid hydrolysis of the substrates p-MPD and 2,4-DPED, respectively. The different shapes of the rate surfactant concentration profiles are related to the different hydrophobicitiesof the substrates (seeDiscussion), which have been observed for many other bimolecular reaction profiles described in the l i t e r t ~ r e . ' , ~ . ' ~The , ~ ' .curves ~ ~ exhibit an observed rate constant maximum (k,mbm)at about 0.060and 0.032 M of NaDP for p-MPD and 2,4-DPED, respectively. The maximum rate constants for both substrates in NaDP are about
IO
15
IO2[ NaDDP] ,M
Figw 3. Plot of observed rate constants for the acid-catalyzed hydrolysis of p-MPD in aqueous solutions of NaDDP at 50 OC.
R3
p-MPD: R, = R2 = H; R3 = OCH3 2,4-DPED: Rl = C2H5;& = R3 = OCH3
15.0
I02[NaDP], M
50 -
-
'". 40E
2
30-
n
0-
20 -
I o lot
01
0
I
5
10
15
IO2[NoDDP]. M
Figure 4. Plot of observed rate constants for the acid-catalyzed hydrolysis of p-MPD in aqueous solutions of NaDDP at pH constant, pH = 5.48 at 50 OC.
16-fold larger than those in water at 50 OC (kob,westimated from kob,w= k2,,[H+Iw) using values of [H'], calculated from the experimental pH in Table I1 at the appropriate surfactant concentration. The observed catalytic factors are similar to those observed in the presence of sodium dodecyl sulfate (SDS), kr!m/kob,w = 10 at 25 oC.lla Figure 3 shows the effect of NaDDP upon kobqm for thepMPD acid hydrolysis. The shape of the rate constantsurfactant concentration profile is very similar to that observed in Figure 1, the is displaced to a lower surfactant only difference being that concentration due to the smaller cmc of the NaDDP. When the pH in the aqueous phase is maintained at 5.48 by addition of small aliquots of HCl(O.1 M), Figure 4, a different behavior in NaDDP solutions is observed. The rate constant increases toward a plateau at high concentrations of surfactant. The observed plateau is consistent with the low binding constant of the substrate and in PPIE terms corresponds to a balance between the expected rate increase with greater substrate binding and the inhibition due to
kzm
1466 The Journal of Physical Chemistry, Vol. 96, No. 3, 1992 TABLE II: DHValues of Micellar NaDP and NaDDP at 50 OC” 102[SURF],
M
NaDP
1.2 1.8 2.4 2.5 2.8 3.0 3.1 3.2 3.5 3.6 4.0 4.2 4.8 5.0 5.4 6.0 7.0 7.2 8.0 8.4 9.0 9.6 10.0 10.8 12.0 13.2 14.4 15.0
4.70 4.78 4.84
NaDPb
PH NaDDF
4.96 5.06
5.20
5.22
5.36
5.38 5.50
5.48 5.56
5.58
5.60
5.43 5.56 5.70
5.28 5.30 5.33
5.79
5.35
5.90
5.37
5.96 6.02
5.40
6.06 6.09
5.46
6.14
5.50
parameters k2,s, M-l s-I cmc
K, kz,,, M-I
s-’
2,4-DPEDa
substrate p-MPD no buffer NaDP NaDDP
12.7 0.027
194 0.036
27 14.3
4 652
194 0.0114b 0.0098‘ 4 854b 784c
in buffer NaDDP 194 0.0035 4 790
“In presence of NaDP micelles. bFrom fit of the data in Figure 3, pH of solution variable. CFromFigure 4, constant pH 5.48.
For ionic micelles, the PPIE model assumes that there is competition between bound counterions and other counterions in the bulk phase. For the acid hydrolysis of p M P D and 2,CDPED in the presence of monoalkyl phosphate surfactants, the exchange of sodium and hydrogen ions by the micelle is described by an ion-exchange constant KH/Nareq 2, where the subscripts m and
6.16 5.56 6.19 5.69
5.65
5.73
5.65
5.76
5.65
6.21 6.22 6.23 6.24 6.26
” PH corresponding to that of data in Figure
5.63 5.65 5.67 5.68
1. pH corresponding to that of data in Figure 2. C p Hcorresponding to that of data in Figure 3. “pH corresponding to that of data in Figure 5.
-6 0
TABLE III: Panmeters Used To Fit the Experimental Results
NaDDPd
4.50 4.74 4.90 4.94 5.05
Ruzza et al.
I
w denote micellar and aqueous phases, respectively. Dioxolane hydrolysis is a bimolecular reaction, and to estimate the second-order rate constant in the NaDP and NaDDP micellar phases, k, the data in Figures 1-5 were treated using eq 3, where
kob,m = [H+lw[(k2,m/ii)KsKH/Na([Na+lm/[Na+lw)+ k2,wl (3) (1 + CdKs) Pis the partial molar volume of the monomer in the micelle; C, is the stoichiometric concentration of surfactants forming micelles defined by c d = C, - cmc, where C, and cmc are the total surfactant and critical micelle concentrations,respectively. The sodium ion concentrations in the micellar ([Na+],) and aqueous ([Na+],) phases were initially calculated on the basis of ion-exchange (eq 2) following the procedure suggested by Quina and Chaimovich2and assuming that the degree of dissociation, a,was constant under all conditions, Le., the total concentration of ions in the micellar surface is constant. To treat the experimental data with eq 3, reported values of (Y = 0.3, P = 0.25, and KH/Na = 1 were used in all c a ~ e s ,and ~~,~~ best fits (represented by the solid lines in Figures 1-5) were contained by minimizing the weighted standard deviations. The hydrolyses of p-MPD and 2,4-DPED in the presence of NaDP (Figures 1 and 2) were best fitted with values of k2,, of 652 and 14.3 M-I s-I, respectively (Table 111). Also included in Table I11 are the K,, k2,,, and cmc values used in fitting the data. The & value for 2,CDPED is about 7-fold greater than that of pMPD, which is consistent with the greater hydrophobicity of 2,4-DPED and which is reflected in the shape of the kinetic profiles. The Ks values are relatively low when compared with parameters obtained at 25 O C ; however, the shape of the kinetic profde r e q h such a small K,value at 50 O C . For p-MPD, an increase in K, from 4 to 50 results in an increase in the standard deviation from 5.5 to 454, a result which indicates that the tit is sensitive to K, values and therefore enforces the use of low K,values for both substrates. A modest rate enhancement, k&kzw = 3, is observed for p-MPD, but 2,4-DPED shows values of kz,mand kz,wwhich are essentially identical. This is different from previous results of micellar effects on acetal and ketal hydrolyses in which the ratios are 0.5 or 1ess.11a*22 The slightly smaller value of the cmc, kinetically detected in the presence of the 2,CDPED substrate, may indicate that micelle formation is promoted by the substrate. This type of effect is not surprising, since it is well known that organic additives reduce the cmc and affect other physicochemical parameters of ionic micelle^.^^-^' Lianos, P.; Lang, J.; a n a , R. J . Phys. Chem. 1984.88, 819. ( 3 6 ) Almgren, M.;Swarup, S . J . Phys. Chem. 1982,86, 4212. (35)
The Journal of Physical Chemistry, Vol. 96, No. 3, 1992 1467
Surfactant Effects on the Acid Hydrolysis of Dioxolanes When NaDDP was used as catalyst in the acid hydrolysis of p-MPD, similar parameters result from treatment of the data (Table 111) despite the fact that different kinetic profiles were obtained under different experimental conditions. The cmc value decreased from 0.0114 M when the pH was variable to 0.0098 M when the pH of the solution was maintained at 5.48. A cmc of 0.0035 M was obtained for the same system in the presence of the succinic acid/sodium succinate buffer. These observed decreases in cmc are consistent with the well-documented salt e f f e c t ~ . l , ~The ~ *values ~ ~ of K, for p-MPD are low and identical in all cases, showing that slight pH variations and addition of up to 20 mM salt do not significantly affect substrate incorporation into the micelle. This is reasonable becausep-MPD does not suffer significant protonation in the pH range under study. The k2,m values obtained for the three different experiments give an average value of 809 f 30 M-' s-I, which is an excellent agreement considering the variation in the experimental conditions (Figures 3-5). The hydrogen ion concentration in the micellar phase, under our experimental conditions, is much less than the concentration of bound counterions, Le., [H+],
