Langmuir 1997, 13, 4079-4084
4079
Surface and Aggregation Properties of N-(4-Nitrophenyl)perfluorononanamide Alejandro M. Granados,† Gerardo D. Fidelio,‡ and Rita H. de Rossi*,† Departamento de Quı´mica Orga´ nica-INFIQC and Departamento de Quı´mica Biolo´ gica-CIQUIBIC, Facultad de Ciencias Quı´micas, Universidad Nacional de Co´ rdoba, Ciudad Universitaria, Agencia Postal 4, Casilla de Correo 61, 5000 Co´ rdoba, Argentina Received March 3, 1997. In Final Form: April 18, 1997X It was determined that N-(4-nitrophenyl)perfluorononanamide (1) forms a stable monolayer which suffers a liquid expanded (LE) f liquid condensed (LC) phase transition upon compression. The surface pressure of the transition was linearly dependent on the temperature indicating a diffuse first-order transition process. The transition was ascribed to rotation of the C-N bond of the amide. In the LE state, 1 has an “L” conformation and in the LC state the phenyl group is aligned with the perfluoroalkyl chain. The pK of 1 is about 2 units higher than the value expected in the water solution. The surface potential of 1 at the air-water interface is negative at all measured pH values, indicating that the molecule on the surface can be considered as a vector having its air end negative. In solution, 1 is in an aggregated form even at 2 × 10-6 M concentration and these aggregates are destroyed in solutions containing 0.01 M of sodium dodecyl sulfate (SDS) or 7 × 10-3 M of perfluorononanoic acid 2. The kinetics of deaggregation occurs with two characteristic times with values of τ1 74.6 and 21.7 s and τ2 294 and 83 s for SDS and 2, respectively, at 25 °C.
Introduction The physicochemical properties of surface active fluorocarbon derivatives have attracted considerable interest in the latest years. These substances are extremely surface active and the critical micelle concentration (cmc) values of the totally fluorinated surfactants are considerably lower than those of the corresponding hydrocarbon compounds.1,2 Previous work dealing with amphiphilic fluorocarbons has established that their cmc depends on the type and concentration of counterions, temperature, and other factors that influence the self-aggregation properties of these surfactants.3 Recently, it has been reported that replacing the hydrocarbon chains of glycerophosphatidylcholine by perfluorocarbons confers a higher stability to liposomes.4 In this connection, was found that amphiphiles with only one perfluorocarbon chain of eight to nine carbons can self-aggregate into uni- or multilamellar vesicles and micelles with different degree of eccentricities.5 Besides, supramolecular organizations were obtained with single-tailed glycolipids while single chain surfactants only form bilayers when special constrains are present.6-10 Enhanced stability of these aqueous aggregates and the chemical inertness of perfluorocarbon compounds (biocompatibility) have been explored to design oxygen carriers in artificial blood.11 †
Departamento de Quı´mica Orga´nica-INFIQC. ‡ Departamento de Quı´mica Biolo ´ gica-CIQUIBIC. X Abstract published in Advance ACS Abstracts, June 1, 1997. (1) Kunieda, H.; Shinoda, K. J. Phys. Chem. 1976, 22, 2468. (2) Ravey, J. C.; Stebe, M. J. Colloid Surf. A 1994, 84, 11-31. (3) Shinoda, K.; Hato, M; Hayashi, T. J. Phys. Chem. 1972, 76, 909. (4) Santaella, C.; Vierling, P.; Riess, J. G. Angew. Chem., Int. Ed. Engl. 1991, 30, 567. (5) Krafft, M. P; Giulieri, F; Riess, J. Angew. Chem., Int. Ed. Engl. 1993, 32, 741. (6) Zarif, L.; Gulikkrzywicki, T.; Riess, J. G.; Pucci, B.; Guedj, C.; Pavia, A. A. Colloid Surf. A 1994, 84, 107. (7) Riess, J. G. Colloid Surf. A 1994, 84, 33. (8) Myrtil, E.; Zarif, L.; Greiner, J.; Riess, J. G.; Pucci, B.; Pavia, A. A. Macromol. Chem. Phys. 1994, 195, 1289. (9) Riess, J. G. Artif. Cells, Blood Substitutes, Immobilization Biothechnol. 1994, 22, 215. (10) Riess, J. G. Biomater., Artif. Cells, Immobilization Biotechnol. 1992, 20, 183.
S0743-7463(97)00234-5 CCC: $14.00
The Langmuir trough has been used to study the surface behavior of amphiphilic perfluorocarbons and related compounds.12-14 Particularly, Yoshio et al.14 have reported that the ability of perfluoroalkylazobenzene derivatives to form insoluble monolayer depends on the chain length of the hydrophobic tail. In previous work,15 we found evidence that N-(4nitrophenyl)perfluorononanamide (1) forms aggregates even at concentrations as low as 2 × 10-6 M and the growing interest in the physicochemical properties of perfluorocarbon aggregates made up of relatively simple amphiphilic molecules prompted us to study the surface properties of 1 at the aqueous-air interface at different temperatures and subphase pHs. We also correlate the surface properties with those measured in bulk conditions. We found that compound 1 forms insoluble monolayer at air-water interface showing a liquid expanded to liquid condensed (LE f LC) transition. Some of the properties of 1 measured at the air-water interface are different from those determined in the bulk due to the different array of the molecules in the two environments. Besides in an aqueous solution of 1, the incorporation of perfluorononanoic acid 2 or sodium dodecyl sulfate (SDS), destroys the aggregates of 1 at a rate remarkably slower than that corresponding to the formation and dissolution of aqueous detergent solutions which lies in the range of 106- 109 s-1.16,17 Results Surface Properties of Perfluorocarbon Derivatives 1 and 2. Compound 1 forms insoluble monomolecular films on a 145 mM NaCl-air interface. The stability of the film was confirmed by the reproducibility of the surface parameters under successive compression and expansion cycles indicating that no desorption of 1 (11) Gross, U.; Papke, G.; Rudige, S. J. Fluorine Chem. 1993, 61, 11. (12) Menger, F. M.; Littau, C. A. J. Am. Chem. Soc. 1993, 115, 10083. (13) Menger, F. M.; Wood, M. G., Jr.; Richardson, Q. Z.; Elrington, A. R.; Sherrod, M. J. J. Am. Chem. Soc. 1988, 110, 6797. (14) Yoshino, N.; Kitamura, M.; Setro, T.; Shibata, Y.; Abe, M.; Ogino, K. Bull. Chem. Soc. Jpn. 1992, 65, 2141-2144. (15) Granados, A.; de Rossi, R. H. J. Am. Chem. Soc. 1995, 117, 3690. (16) Muller, N. J. Phys. Chem. 1972, 76, 3017. (17) Frindi, M.; Michels, B.; Zana, R. J. Phys. Chem. 1992, 96, 6095.
© 1997 American Chemical Society
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Figure 3. Schematic representation of compound 1 in the airwater interface.
Figure 1. Isotherms of 1 at different temperatures: (A) ‚‚‚, 7 °C; - - -, 11 °C; -, 24 °C; (B) - - -, 5 °C; -‚-, 13 °C; ‚‚‚, 17 °C; -, 26 °C. Subphase was 0.145 M NaCl in water.
Figure 4. Change of the surface potential of the monolayer of 1 as a function of pH.
Figure 2. Plot of the phase transition pressure vs temperature.
takes place during compression. When the isotherm of 1 was run at room temperature (24 °C), it showed a clear discontinuity at around 5 mN‚m-1 (Figure 1) revealing the possibility of bidimensional liquid expanded (LE) to liquid condensed (LC) phase transition (LE f LC), similar to that observed for glycosphingolipids18 and phospholipids.19 If this discontinuity is due to a LE f LC transition, the surface pressure point of the transition should depend on the temperature of the system.18 This was confirmed by doing the isotherms of 1 at different temperatures between 5 and 32 °C (Figure 1). Figure 2 shows that there is a linear dependence of the surface pressure at which the LE f LC transition initiates (πt) with the temperature. This result is indicating that the discontinuities found in the π-area isotherms can be considered as a diffuse first-order transition process.18,20 (18) Fidelio, G. D.; Maggio, B.; Cumar, F. A. Biochim. Biophys. Acta 1986, 854, 231. (19) Fidelio, G. D.; Maggio, B.; Cumar, F. A. Biochim. Biophys. Acta 1986, 49. (20) Georgallos, A.; Pink, D. A. J. Colloid Interface Sci. 1982, 89, 107.
The phase transition was not observed at temperatures above 30 °C or at a pH higher than 11 at room temperature. The study was done at 7 < pH < 11 because at pH higher than 11 no phase transition was observed. The molecular area at the collapse pressure (the maximal lateral pressure that can be supported before the film disrupts its monomolecular array) at room temperature, 24 °C, is about 0.39 nm2. The corresponding molecular area at πt is 0.64 nm2. Comparison of these experimental values with those obtained with space-filling CPK molecular models shows that the phase transition is compatible with a change in the molecular conformation involving a rotation of the C-N bond of the amide group. In the LE state, 1 acquires an “L” conformation with the perfluorocarbon chain perpendicular to the interface and the p-nitrophenyl group parallel to the interface. In the LC state this group rotates almost 90 ° and it aligns to the main axis perpendicular to interface (Figure 3). We measured the surface potential (∆V) of the monolayer of 1 at the air-water interface as a function of the pH in order to determine the value of the pKa of the -NH group (Figure 4). According to Gaines21 the difference of electrostatic potential between the air and the aqueous phase ∆V is given by eq 1 (21) Gaines, G. L. In Interscience Monographs in Physical Chemistry: Insoluble Monolayers at Liquid-Gas Interfaces; Prigogine, I., Ed.; John Wiley: New York, 1966.
Surface Active Fluorocarbon Derivatives
∆V ) (0.38µ⊥/A) + Ψo
Langmuir, Vol. 13, No. 15, 1997 4081
(1)
Table 1. Characteristic Times for the Incorporation of 1 in 2 Aggregates
where A is the molecular area in nm2, µ⊥ is the overall dipole moment perpendicular to the interface, and Ψo is the Guy-Chapman surface potential. The value of ∆V results from several contributions to the overall molecular moment µ⊥, as indicated in eq 2.
µ⊥ ) µw + µp + µh
(2)
In this equation µw, µp, and µh represent the contribution from the oriented water molecules around the polar head group, the fundamental dipoles in the polar head group, and the hydrophobic chain dipoles of the amphiphile, respectively. As the titration takes place, ∆V changes due to the variation of at least three of the parameters of eq 1: µw, µp, and Ψo. Since ∆V depends, among other factors, on the state of protonation of the ionizable groups, it can be used to estimate an apparent pKa of about 10.4 (Figure 4), in agreement with that obtained in bulk solution (see below). It can be seen from Figure 4 that the surface potentials at any pH are negative, which indicates that the overall dipole moment of the molecule can be considered as a vector having its air end negative.21 On the other hand, amphiphilic hydrocarbons such as fatty acids and phospholipid21 and glicoesphingolipids22 give positive values. Neither the equivalent hydrocarbon amide, i.e. pnitrophenylnonanamide (3), nor compound 2 formed insoluble monolayer at the air-water interface. Both compounds showed a strong liquid expanded character and become soluble upon compression but it could be determined that for 3 the tendency in surface potentials is positive. The behavior of 2 and 3 as compared with that of 1 indicates that both the perfluorocarbon chain and the p-nitrophenyl group are needed for interfacial stability. Bulk Thermotropic Behavior of 1 and 2. It has been reported that for phospholipids as well as for gycosphingolipids, the temperature dependence of the monolayer that undergoes bidimensional phase transition can be used to estimate the equivalent bulk transition temperature.18,23,24 This approach assumes that the molecules in the aqueous dispersion have intermolecular interactions similar to those acquired in the monomolecular array.18,23 If a similar comparison is made for perfluorocarbons, the bulk phase transitions should be around 30 °C considering the data shown in Figures 1 and 2. Unexpectedly, using high sensitivity differential scanning calorimeter only one cooperative heat absorption peak was observed at 109 °C in a pressurized aqueous solution of 1. This temperature is close to the melting point of pure 1, namely, 119-120 °C.15 These results might indicate that compound 1 in water aggregates with a completely different long range molecular organization from that in the monolayer. Compound 2 showed a cooperative heat adsorption peak at 45 °C. This temperature is close to the Krafft point25,26 and lower than the melting point, namely, 77 °C.27 These (22) Maggio, B.; Cumar, F. A.; Caputto, R. Biochem. J. 1978, 171, 559. (23) Maggio, B.; Fidelio, G. D.; Cumar, F. A.; Yu, R. K. Chem. Phys. Lipids 1986, 42, 4963. (24) Phillips, M. C.; Chapman, D. Biochim. Biophys. Acta 1968, 163, 301. (25) The Krafft point is the temperature at which the solubility of a surfactant increases sharply as a consequence of the formation of micelles. See ref 26. (26) Mayer, Drew Surfactant Science and Technology, 2nd ed.; VCH Publishers, Inc. Dearfield Beach, FL, 1992; p 83. (27) La Mesa, C.; Sesta, B. J. Phys. Chem. 1987, 91, 1450.
b
T, °C
τ-11, 10-2 s-1 a
τ-12, 10-3 s-1 a
25 35 45 50
1.34 (4.6)b 2.35 3.91 17.7
3.4 (12)b 6.7 9.2 31.5
a In all cases the standard deviation of the fit is less than 2%. Values in brackets are for solutions of SDS (0.01 M).
data and the inability of 2 to form insoluble monolayer are indicating that the extent of hydration of compound 2 is higher than that of 1 in aqueous bulk dispersion Spectrophotometric Determinations. The absorbance of 2 was measured as a function of its concentration in water at 25 °C. The increment in absorbance was linear up to 2.4 × 10-3 M (plot not shown). Above this concentration, the slope increased but the tendency is also linear. This change in slope may be associated with the formation of submicellar aggregates since the temperature is below the reported Krafft point for this compound, namely, 43.8 °C27 and 48.3 °C.28 Similar changes in slope were observed for perfluoroheptanoic acid and the change in slope at 230 nm which takes place at 0.0328 M is considered the cmc.29 The dependence of the absorbance with the concentration of compound 1 was linear in the range tested (from 10-6 to 2 × 10-5 M) at 25 °C and pH 10 although it is in the aggregated form. This behavior was similar at the three wavelengths used: 270 nm (λmax of the aggregate), 300 nm (λmax of compound 1 in the monomeric form), and 334 nm (λmax of compound 1 in the deprotonated form). The pH 10 was used because 1 has a higher solubility at this pH than at lower pH. The λmax at 270 nm is attributed to the aggregated form since the normal maximum absorption found for shorter chain compounds (amides with (CF2)n with n e 5) is 300 nm.15 Kinetic Measurements. When a solution of 1 in acetonitrile was added to a solution of 2, 7 × 10-3 M, at pH 3 the maximum absorption shifted from 270 to 300 nm, i.e., the final spectrum is that of the monomeric 1.30 There are two isosbestic points one at 284 and another at 290 nm, the first appears at shorter and the latter at longer times. These results indicate that the presence of 2 destroys the aggregates of 1. The change in absorbance with time was measured at 300 nm and the values were fitted to a double exponential function (see Experimental Section) from which two characteristic times τ1 and τ2 can be extracted. Similar results are obtained in the presence of 0.01 M SDS. On the other hand, a cationic detergent cetyltrimethylammonium bromide (CTAB) at 0.01 M has no effect on the spectrum of 1 even after 24 h. The kinetic was measured at temperatures between 25 and 50 °C and the data are collected in Table 1. Arrhenius type plots (not shown) for τ-11 and τ-12 are curved and show higher slope in the range 45-50 °C than in the range 45-25 °C. The slope of the plot of log τ-11 vs the reciprocal of the absolute temperature changes from -2.21 to -13.1 K and that of τ-12 changes from -2.1 to -10.7 K. Discussion Some of the surface parameters of compound 1 measured at the air-water interface are in agreement with those (28) Kunieda, H.; Shinoda, K. J. Phys. Chem. 1976, 80, 2468. (29) Mukerjee, P.; Gumkowski, M. J.; Chan, C. C.; Sharma, R. J. Phys. Chem. 1990, 94, 8823. (30) As the spectrum of 1 dissolved in acetonitrile does not change with concentration in the range 10-5 to 10-3 M, it probably is in its monomeric form in this solvent.
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obtained for the aggregate in the bulk solution but others are not. The apparent pKaapp measured in both forms and with different approaches has essentially the same value, i.e., 10.2, which is nearly 3 units higher than that reported for similar amides with shorter perfluorinated alkyl chain in water solution.15 It is well-known that the intrinsic pKa value for charged groups at the interface can be influenced by several surface parameters not present in the monomeric soluble form of the amphiphile. The interfacial intrinsic pKai of a hydrocarbon phospholipid in a monolayer or an aggregated form is given by eq 331
pKai ) pKaapp + ∆ pKapol + eψ/2.303kT
(3)
where ∆pKapol is the shift in the acid-base equilibrium of neutral acid due to the different polarity of the interface compared to the bulk environment of a monomer. The term ψ is the electrostatic surface potential which determines the actual proton concentration at the interface. The value of ψ is around -140 mV for the fully charged monolayer of compound 1 at collapse pressure. If ∆pKapol is similar to the value reported for phospholipids, namely, 1,31 pKai should be around 9. This value is 2 units higher than those reported for compounds that are soluble with a shorter chain.15 These results may indicate that the NH group is involved in the stabilization of aggregated forms both in the monolayer and in bulk for some specific intermolecular interaction. The idea is further supported by the fact that at pH above 11 (fully charged species) the aggregate increases its solubility as monomeric anion (λmax shift to 300 nm) and also at this pH the monolayer increases its liquid character, in agreement with a lower intermolecular interaction. The neutralization of the surface charges of micelles of perfluoro acids and the resultant reduction of electrostatic destabilization of micelles were considered to be responsible for an increased stability of fluorocarbon acid micelles as compared with their salts.29 The cmc’s of perfluorooctanoic acid and its sodium salt are 0.0096 and 0.0306 M, respectively, at 25 °C.29 In contrast, the cmc of dodecylsulfonic acid is only 1.15 times lower than that of its sodium salt.32 The linear dependence of the initial surface pressure of the bidimensional phase transition πt with temperature, allows the assumption of a diffuse first-order transition; therefore the process can be described by the twodimensional Clayperon equation (eq 4).19,24
dπt Qc ) dT T(Ae - Ac)
(4)
In eq 4 πt is the surface pressure at which LE f LC transition begins, Qc is the latent heat associated to the transition, T is the temperature of the subphase, Ae is the molecular area at πt, and Ac is the molecular area at the beginning of the LC portion of the isotherm and is obtained by extrapolation of the LC portion down to πt. The dπt/dT is obtained from the slope of the plot of πt vs temperature (Figure 2). The dπt/dT is constant with temperature and gives an idea of the entropy changes associated with the transition or the entropy changes per unit of area for the condensation process; this can also be concluded from eq 4 written as eq 5. (31) Tocanne, J. F.; Teissie, J. Biochim. Biophys. Acta 1990, 1031, 111. (32) Mukerjee, P.; Mysels, K. J. Critical Micelle Concentration of Aqueous Surfactant Systems. Natl Stand. Ref. Data Ser. (U. S.> Natl. Bur. Stand.) 1971, NSRDS-NBS 36.
dπt ∆St ) dT ∆A
(5)
The value of dπt/dT for compound 1 is 0.23 mN‚m-1‚°C-1 and much lower than the values reported for dimyristoylphosphatidylcholine and dipalmitoylphosphatidylcholine, which are 1.5 mN‚m-1‚°C-1 or 1.7 mN‚m-1, respectively.19 Both hydrocarbon phospholipids have comparable surface areas and undergo two-dimensional phase transitions at similar temperature. These data are indicating that the entropy changes associated with fluorocarbon transition in the monolayer are lower than the equivalent values for phospholipid. This can be attributed both to higher difficulties in the trans-gauche isomer formation for the fluorocarbon compared to hydrocarbon chain and to a lower capacity of compound 1 to hydrate its hydrophilic head group in the liquidexpanded region at neutral pH. The latter conclusion is also in agreement with the high tendency of 1 to aggregate at very low concentration. The greater van der Waals radius of fluorine relative to hydrogen and the longer C-F than C-H distances (by ca. 0.2 Å)33 lead to a larger crosssection area for fluorocarbon than for hydrocarbon chains and the energy difference between the trans and gauche conformation is also larger for fluorocarbon than for hydrocarbons.34 For hydrocarbon phospholipids as well as for glycosphingolipids a good correlation between monolayer and the thermotropic behavior in bulk aggregates was found. The corresponding bulk phase transition temperature can be estimated from the surface behavior of the isotherms with the temperature.18,23 The correlation is possible since in these aggregates (micelles, cylindrical micelles, or bilayer vesicles) the interaction of the hydrophobic hydrocarbon chains in the bulk phase is not substantially different from that found for the monolayer. In the case of the perfluorocarbon compound 1, the expected bulk phase transition estimated from the monolayer data is around 30 °C. However, this temperature measured by differential scanning calorimeter is 109 °C, close to the melting point. This is indicating that hydration is limited and the long range organization in bulk is rather different from that of the monolayer array. In the monolayer array, 1 has the most hydrophilic portion protruding toward the aqueous phase whereas the perfluorocarbon portion lies in the more hydrophobic region of the interface giving an overall dipole with its negative end toward the air. We suggest that the aggregates in the bulk are in a head-tail arrangement as shown in Figure 5. This organization avoids unfavorable dipolar interactions and could account for the high thermal stability found for the bulk aqueous aggregate. Compound 2 does not form a stable monolayer at the air-water interface and the amphiphile is solubilized into the subphase upon film compression; this occurs also at pH ) 1 where the compound should be fully protonated. This indicates that the enhanced hydrophobicity given by the perfluorocarbon chain alone is not enough to confer to the molecule an appropriate surface stability. The replacement of the OH group of the acid 2 by the p-nitrophenylamino group gives to compound 1 not only further hydrophobicity but also the possibility of attractive intermolecular interaction probably throughout the hydrogen bond between the NH group and the CdO of the adjacent molecule. As the pH of the subphase increases (33) The C-F and C-H bond distances were computed using the MM2 method of Chemoffice 3D software. It also computed the stretching and bending energy for 1 and 3. These values, in kcal/mol units, are 13.8 and 7.46 for 1 and 0.8 and 2.57 for 3. (34) Fontell, D.; Lindman, B. J. Phys. Chem. 1983, 87, 3289.
Surface Active Fluorocarbon Derivatives
Langmuir, Vol. 13, No. 15, 1997 4083
from emulsion droplets to the microsomal membranes.37 Under our experimental conditions, the aggregation of 1 occurs after the mixing of the stock solution in acetonitrile to water and before the first UV spectrum can be taken (∼1 s).30 The overall process can be described as in eqs 6, 7, and 8 where 1M and 1A represent compound 1 in the monomeric and aggregated forms, respectively, and (1M.Micelle)I and (1M.Micelle)II represent two different types of association of 1 with the micelles of SDS or compound 2
Figure 5. Schematic representation of an aqueous aggregate of 1.
and the titration of the proton takes place, the liquid character of the monolayer of compound 1 increases. Compound 2 has an endothermic absorption in aqueous solution centered at 45 °C. This transition can be related to the heat absorption at the Krafft point reported in the literature.29 Before discussing the effect of the addition of SDS and 2 to solutions of 1, it is necessary to make some considerations regarding the state of the surfactants. According to the phase diagram reported,34 2 is in the two phase region between 25 and 45 °C: isotropic solution and liquid crystal phase which appears at concentration around 2.4 × 10-3 M as determined by the change in absorbance vs concentration. Above the Krafft point, around 45 °C, the micellar phase is formed. The critical micellar concentration of 2 at 60 °C is 2.8 × 10-3 M. On the other hand considering that the Krafft point of SDS is 16 °C and that the cmc at 25 °C is 0.008 M, a 0.01 M solution of SDS is in its micellar phase. The solubility of compound 1 increases in the presence of 2 or SDS as suggested by the spectrophotometric data. The kinetics of the formation of monomeric 1 in the presence of aggregates of 2 or SDS micelles can be described by a two exponential equation with two characteristic times τ1 and τ2 indicating that the process occurs in at least two steps. It is known that the incorporation of dyes into SDS micelles is given by a two-step process.35 In this case the faster event, which takes place in the microsecond time scale, is associated with the adsorption of the dye on the micelle surface. The slower event, in the millisecond time scale, is attributed to the incorporation of the dye to the micelle.36 The two kinetic processes of our system occur at much slower rate, the time scale is seconds for the first and minutes for the second (Table 1). Similar slow processes were reported for the complex formation between microsomal cytochrome P450 and perfluorocarbons and it was suggested that the limiting stage in the whole system is the diffusion of the perfluorocarbon (35) Fendler, J. H. Membrane Mimetic Chemistry; John Wiley & Sons: New York, 1982; p 31. (36) Takeda, K.; Tatsumoto, N.; Yasunaga, T. J. Colloid Interface Sci. 1973, 495.
1M h 1A
(6)
1M + Micelle h (1M.Micelle)I
(7)
(1M.Micelle)I h (1M.Micelle)II
(8)
We suggest that the two processes that we measured are associated with equilibriums 7 and 8 since equilibrium 6 must be very fast because in the absence of SDS or 2 we detected only 1M. The deaggregation of 1A must also be fast because only the spectrum of 1M was observed after the addition of β-cyclodextrin to a solution of 1A.15 Besides, the concentration of 1M must be very small. It was reported that perfluoroheptanoic acid forms large aggregates and that the rate of exchange of monomer and aggregate decreases by about 2 orders of magnitude when it is compared with normal rate of hydrocarbon surfactants; for instance, the mean lifetime for exchange of perfluoroheptanoic acid between monomers and micelles is 1.2 × 10-4 s and for perfluorheptanoil polyethoxylated amide is 0.13 s.38 The latter value is about 3 orders of magnitude faster than τ1. The fact that the Arrhenius plot for τ-11 and τ-12 shows a brake at 45 °C may be a manifestation of the phase change of 2 taken place at this temperature as mentioned above. The values of τ1 and τ2 are considerably more sensitive to the change in temperature when 2 is in the micellar phase (above the Krafft point) than when it is in the liquid crystalline phase, but we do not think it is easy to explain these results with the data at hand considering that τ-11 and τ-12 represent a combination of rate constants associated with the strongly coupled process described by eqs 6-8. Conclusions The perfluoro compounds studied behave in the airwater interface as phospholipids of 14-16 carbon atoms. But there are some remarkable differences. The surface potential of the monolayer is negative at all pH, i.e., for the neutral fluorocarbon as well as for the deprotonated compound. This result is opposite to what was found for the hydrocarbon amphiphiles. The data of the physical state in the monolayer do not correspond with that in the bulk solution which may indicate an increase in the intermolecular interaction in solution probably as a consequence of head to tail interactions which maximize the dipolar interactions in the aggregate and decrease the molecular surface exposed to water. Experimental Section Compound 1 was synthesized and purified as reported previously.15 The measurements of the surface properties of the monomolecular films of all the compounds assayed were done in a home(37) Gross, K.; Obraztsov, V. V.; Makarov, K. N.; Radeck, W.; Ru¨diger, S. J. Fluorine Chem. 1993, 63, 101. (38) Guo, W.; Brown, T. A.; Fung, NB. M. J. Phys. Chem. 1991, 95, 1829.
4084 Langmuir, Vol. 13, No. 15, 1997 made monolayer apparatus described in detail previously.18,22 The surface pressure was recorded using the Wilhenmy method by using a platinized-platinum dipping wire of known diameter suspended from an LM 600 Beckman electrobalance;22 the surface potential was measured with a Beckman Zeromatic SS-3 pH meter connected to an air-ionizing electrode of 241Am suspended at 5-7 mm from the air-water interface and to a calomel reference electrode connected to the subphase through a salt bridge.22 The surface pressure and surface potential were recorded in a dual channel x-y recorder and digitized with an appropriate data acquisition system. The spreading solution was 1 (5 × 10-4 M) dissolved in chloroform/methanol (2:1, by vol.; HPLC grade, Sintorgan, Buenos Aires, Argentina). The stock solution (20 µL) was directly spread over the 145 mM NaCl subphase at the specified pH. The desired pH was prepared with HCl or NaOH (1 M) and controlled by an Orion 720 pH meter. The monolayer was compressed isothermally at constant velocity (7.1 cm‚min-1) and the limiting molecular area was defined as the molecular area at the collapse pressure.18,21,22 The temperature was maintained with a Haake F3C circulating bath. The temperature of the trough was recorded with a precision better than 0.1 °C with a digital thermometer probe immersed directly into the subphase. All isotherms were done at least in triplicate. The reproducibility was better than (0.1 mN‚m-1 for surface pressure, (5 mV for surface potential, and (0.02 nm2 for molecular area. In order to get reproducible results, solutions of 2 were prepared 48 h before use. Thermograms were recorded in a Microcal MC2D differential scanning calorimeter. Aqueous solutions of 1 mg/mL of 1 or 2 were extensively sonicated and degassed before loading the calorimetric cell. The scan rate was 1 °C/min. The spectrophotometric measurements were done with a Shimadzu UV 260 or a Shimadzu 2101 PC. For the kinetic runs,
Granados et al. 10 µL of a stock solution of 1 in acetonitrile were injected into a 1 cm optical pass length quartz cuvette containing a water solution of 2 (7 × 10-3 M) of SDS (0.01 M) at pH 3. The final concentration of 1 was 10-5 M. The absorbance readings (Abs) were taken each 0.5 s and the values were fitted using an appropriate software to eq 9 where ao is the final absorbance of the system, and a1 and a2 are related to the amplitude of the two kinetic processes.
Abs ) ao + a1 exp(-t/τ1) + a2 exp(-t/τ2)
(9)
The kinetic data were determined by preparing the solution of 2 (7 × 10-3 M) 42 h before measurements, freshly prepared solutions gave no reproducible results as did solutions prepared 42-96 h earlier. The absorbance values as a function of the concentration of 1 were taken in solutions prepared 24 h before; the solutions were thermostated 2 h before measurements. The change in absorbance with time was determined at 300 nm.
Acknowledgment. This research was supported in part by the Consejo Nacional de Investigaciones Cientificas y Te´cnicas, Argentina (CONICET), and the Consejo Provincial de Investigaciones Cientı´ficas y Tecnolo´gicas de Co´rdoba, Argentina (CONICOR), and the University of Cordoba. A. Granados is a grateful recipient of a fellowship from CONICET. We thank Professor Takaaki Sonada, Kiushu University, Japan, for sending us a free sample of nonanoic acid. LA9702345
