Behavior of 1, 1, 1-Trifluoroacetylacetone and 1, 1, 1-Trifluoro-3-(2

The UV−vis absorption spectra of 1,1,1-trifluoracetylacetone and 1,1,1-trifluoro-3-(2-thenoyl)acetone are studied in water and in aqueous micellar s...
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Langmuir 2000, 16, 8438-8446

Behavior of 1,1,1-Trifluoroacetylacetone and 1,1,1-Trifluoro-3-(2-thenoyl)acetone in Aqueous Micellar Solutions of the Cationic Surfactants Tetradecyltrimethylammonium Bromide and Tetradecyltrimethylammonium Chloride Emilia Iglesias Departamento de Quı´mica Fı´sica e E. Q. I., Facultad de Ciencias, Universidad de La Corun˜ a, 15071 La Corun˜ a, Spain Received June 6, 2000. In Final Form: August 7, 2000 The UV-vis absorption spectra of 1,1,1-trifluoracetylacetone and 1,1,1-trifluoro-3-(2-thenoyl)acetone are studied in water and in aqueous micellar solutions of cationic surfactants forming micelles. In strong acid medium, the presence of micelles does not change the spectra of these trifluoro-diketones; however, in aqueous buffered solutions of acetic acid-acetate, addition of surfactant causes the formation of enolate anions. In strong alkaline medium, or in carbonate-bicarbonate buffer solution, the enolate decomposes. The rate of decomposition is reduced strongly by the presence of cationic micelles. The quantitative treatment of spectral changes measured as a function of pH, in both the absence and presence of surfactants, allows us to determine the keto-enol equilibrium constants, as well as the acidity equilibrium constant of both the enol and keto tautomers. Both 1,1,1-trifluoroacetylacetone and 1,1,1-trifluoro-3-(2-thenoyl)acetone are less than 2% enolized in water. The presence of micelles does not increase the enol content; by contrast, a strong increase of their acidity by approximately two pKa units is observed. The enol of these two trifluorodiketones is more acidic than the keto tautomer, a common observed phenomenon when the enol content is very low.

Introduction The trifluoro derivatives of acetylacetone (AcAc) have been the topic of several studies related to their ketoenol equilibrium,1-3 to their reactivity toward electrophilic reagents (in halogenation, nitrosation reactions),4,5 or to their general use as versatile ligands.6,7 These compounds should play an important role in metal complexation, and for that, they are used as extractants in nonpolar solvents, because they form relatively stable chelating complexes with a variety of hydrated metals ions. Most β-diketones, whether solids, pure liquids, in solution, or in the gas phase, are in equilibrium with substantial amounts of cis-enols, a characteristic that makes easier the study of their keto-enol equilibria. The stabilization of cis-enols by strong H-bonds between the OH and the remaining carbonyl group constitutes the main reason for the influence on the enol content of both the nature of the solvent and the chemical substituents adjacent to the carbonyl group. It is a generally observed trend that the enol content of a β-diketone increases in nonpolar and/or aprotic solvents, that is, when the intermolecular H-bonds with the solvent are not possible, (1) Burdet, J. L.; Rogers, M. T. J. Am. Chem. Soc. 1964, 86, 2105. Allen, G.; Dwek, R. A. J. Chem. Soc. B 1966, 162. Basseti, M.; Cerichelli, G.; Floris, B. Tetrahedron 1988, 44, 2997. (2) Wallen, S. L.; Yonker, C. R.; Phelps, C. L.; Wai, C. M. J. Chem. Soc., Faraday Trans. 1997, 93 (14), 2391. (3) Ellinger, M.; Duschner, H.; Starke, K. J. Inorg. Nucl. Chem. 1978, 40, 1063. Bankowska, Z. Spectrochim. Acta 1967, 23A, 505. (4) Reid, J. C.; Calvin, M. J. Am. Chem. Soc. 1950, 72, 2948. (5) Crookes, M. J.; Roy, P.; Williams, D. L. H. J. Chem. Soc., Perkin Trans. 2 1989, 1015. (6) Le, Q. T. H.; Umetani, S.; Suzuki, M.; Matsui, M. J. Chem. Soc., Chem. Commun. 1995, 2271. (7) Komatsu, Y.; Honda, H.; Sekine, T. J. Inorg. Nucl. Chem. 1976, 38, 1861. Sekine, T.; Yumikura, J.; Komatsu, Y. Bull. Chem. Soc. Jpn. 1973, 46, 2356. Sekine, T.; Yomatsu, Y. J. Inorg. Nucl. Chem. 1975, 37, 185.

and, then, there is no competition with intramolecular H-bonds in the enol. (Obviously, this refers to the enols with H-bonding capability; for example, this will not be the case of dimedone.) Nevertheless, the effects of substituents, especially electron-withdrawing substituents, are not unequivocally documented. Starting with AcAc as the reference molecule, one observes that replacing the CH3 group by the phenyl group, Ph, increases the enol content; by contrast, the replacement by an ethoxy group, CH3CH2O, to give a β-keto-ester, causes a decrease in the enol content.8 The substitution of CH3 by CH2X, CHX2, or CX3 (X ) Cl, F) should produce the same effect as the ester function, because they are electron-withdrawing substituents. A difference could be observed in those compounds with fluoro substituents capable of forming H-bonds between the enolic HO and F atoms. The enol content observed in the chloro derivatives of ethylacetoacetate (EAA) varies in the “normal” way; that is, increasing the number of Cl atoms in the CH3 substituent increases also the stability of the keto tautomer.9 With fluoro substituents one can find serious discrepancies in the literature. In the excellent review by Toullec10 on keto-enol equilibria, including the effect of fluorinated substituents, the literature reported on the topic points in the direction that the fluoro substituents adjacent to the carbonyl group cause a large increase of the stability of the enol tautomer. However, Reid and Calvin4 have determined the enolization equilibrium constants of several diketones, including 1,1,1-trifluoroacetylacetone (TFA) and 1,1,1-trifluoro-3-(2-thenoyl)acetone (TTeA), at (8) Moriyasu, M.; Kato, A.; Hashimoto, Y. J. Chem. Soc., Perkin Trans. 2 1986, 515. (9) Bankowska, Z.; Bukowski, P.; Grabowski, B. Rocz. Chem. 1970, 44, 1481. (10) Toullec, J. In The Chemistry of Enols; Rappoport, A., Ed.; John Wiley & Sons: Chichester, England, 1990; pp 323-398.

10.1021/la000800n CCC: $19.00 © 2000 American Chemical Society Published on Web 10/07/2000

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Chart 1

25 °C. They have used the bromine titration method, whose accuracy has been questioned especially for those substrates with high enol content; in fact, their result of the enolization equilibrium constant of benzoylacetone differs substantially from recently reported data.11,12 Nevertheless, the enolization equilibrium constants they determined for both trifluoro derivatives of AcAc compare quite well with those determined in the present work (vide infra), and the results indicate that TFA and TTeA are less than 2% enolized in water. Recently, we used aqueous micellar solutions to determine enolization constants of β-dicarbonyl compounds in water.12 Basically, the newly developed method is based on the properties of micelles to offer a different medium inside the bulk water solvent and on the extremely solventsensitive character of the keto-enol equilibrium of β-diketones. The spectral changes on the UV-vis absorption spectrum of diketones provoked by micelles are analyzed quantitatively to determine the enol content. In the present work, we apply the same method to the case of TFA and TTeA. In Chart 1 we show previously obtained results12,13 along with the present results. However, the novelty of the work is the effect of cationic micelles on increasing the acidity of both TFA and TTeA. The presence of small amounts of cationic micelles enhances formation of enolate anions, even at pH values lower than the pKa of the ketone. On the other hand, the results obtained in the nitrosation of TFA and TTeA are in complete agreement with the low enol content observed for these compounds in water; in fact, we concluded that the enolate is the only reactive species.14 Experimental Section 1,1,1-Trifluoroacetylacetone and 1,1,1-trifluoro-3-(2-thenoyl)acetone were Aldrich products of the highest purity and were used as supplied. The surfactant tetradecyltrimethylammonium bromide (TTABr), a Sigma product of the highest available purity, was used without further purification. Tetradecyltrimethylammonium chloride (TTACl) was prepared through the ion exchange of a solution of TTABr, using an Amberlite IRA-400(Cl) ionexchange resin, followed by elution with distilled water. All other reagents were supplied by Merck and were used as received. All the solutions were prepared with water doubly distilled from permanganate solution. The ketones were dissolved in dioxane. From these stock solutions, the working solution was prepared daily by diluting the appropriate volume of the ketone-dioxane solution in the desired volume of water to give aqueous solutions of TFA in the concentration range (2-4) × 10-3 M (and around 10-3 M in the case of TTeA). The percentage of dioxane in the final sample mixture was always less than 0.2 vol %. (Experiments were also performed by adding the diketone dissolved in dioxane directly to the reaction sample, 40 µL for the reaction in water or 4-10 µL for the reaction in micellar medium in 3 mL total volume. We (11) Bunting, J. W.; Kanter, J. P.; Nelander, R.; Wu, Z. Can. J. Chem. 1995, 73, 1305. (12) Iglesias, E. J. Phys. Chem. 1996, 100, 12592. (13) Iglesias, E. J. Chem. Soc., Perkin Trans. 2 1997, 431. (14) Iglesias, E. Results to be published.

Figure 1. Absorption spectra of TTeA (6.4 × 10-5 M) dissolved in (A) water and in (B) aqueous micellar solutions of 0.022 M TTABr at (1) [H+] ) 0.033 M (HBr), (2) 0.067 M acetic acidacetate buffer of pH 4.15, (3) [OH-] ) 0.033 M, and (4) 0.051 M carbonate-bicarbonate buffer of pH 10.2. observed that the absorbance readings are higher than those observed by following the previous procedure; differences are very noticeable for the absorption band due to enolate, especially in the case of TTeA.) The pH was controlled using buffer solutions of acetic acidacetate, which were prepared with acetic acid and NaOH. pH was measured with a Crison 2001 pH-meter equipped with a GK2401B combined glass electrode and calibrated using commercial buffers of pH 4.01, 7.02, and 9.26 (Crison). The reported pH values refer to the pH of the same sample used to record the spectra. The reported [buffer] refers to the total buffer concentration. Ultraviolet absorption spectra were recorded with a KontronUvikon (Model 941) double-beam spectrophotometer with a thermostated cell holder. Each scan was taken after 15 min (to allow the sample to reach a constant temperature). The reference cell (l ) 1 cm, the same for the sample cell) contained all the reagents except for the ketone. All measurements were performed at 25.0 ( 0.1 °C.

Results Figure 1 shows the spectra of TTeA (6.4 × 10-5 M) dissolved (a) in water and (b) in aqueous micellar solutions of TTABr (0.022 M) under different acid-base conditions: strong acid (HBr), acetic acid-acetate buffer, carbonate-bicarbonate buffer, or alkaline medium (NaOH). In aqueous acid medium, the spectra show a broad absorption band with a peak centered at 266 nm and a shoulder at 292 nm (the spectrum in the presence of 0.033 M HBr matchs that in 0.067 M acetic acid-acetate buffer at pH 4.15. The spectrum in alkaline medium is not different except by a small shift to lower wavelength. In aqueous 0.051 M carbonate-bicarbonate buffer at pH

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Figure 2. Absorption spectra of TFA ([TFA] ) 1.1 mM in water and 0.12 mM in aqueous micellar medium) (A) in acid medium at (1) [H+] ) 0.040 M (HCl) and [TTACl] ) 0.044 M, (2) [H+] ) 0.040 M, (3) 0.068 M acetic acid-acetate buffer of pH 4.15, and (4) 0.068 M acetic acid-acetate buffer of pH 4.06, [TTACl] ) 0.044 M and (B) in basic medium at (1) [OH-] ) 0.033 M, [TTACl] ) 0.044 M, (2) [OH-] ) 0.033 M, (3) 0.051 M carbonatebicarbonate buffer of pH 10.2, [TFA] ) 1.2 × 10-4 M, and (4) 0.051 M carbonate-bicarbonate buffer of pH 10.2, [TTACl] ) 0.044 M.

10.20 a new absorption band maximum is observed at 341 nm. However, when TTeA dissolved in dioxane (0.0406 M) is added directly (5 µL) to the sample mixture (3 mL), the absorbance values are higher and also it is possible to observed the absorption band due to enolate in alkaline medium. The absorbance intensity is even higher in carbonate-bicarbonate buffer, although the decomposition rate of enolate is higher in the buffer solution than in alkaline medium. In aqueous micellar solutions of TTABr (Figure 1b), neither the spectrum in strong acid nor the spectrum in basic media (alkaline or carbonatebicarbonate buffer medium) differed significantly. By contrast, for the spectrum recorded in the presence of acetic acid-acetate buffer, a strong absorption maximum is observed at 341 nm, while the absorption intensity of both the peaks centered at 266 and 292 nm decreases appreciably. Remembering the experimental facts observed with benzoylacetone (BZA) (see below), we might attribute the absorption band at 341 nm to enolate anions of TTeA. The results in Figure 1 indicate that the presence of cationic micelles causes a strong increase of the ionization degree of keto-enol tautomers of TTeA, due to the stabilization by electrostatic attractions of the enolate anions. In fact, anionic micelles of SDS do not affect at all the absorption spectrum of TTeA. Furthermore, in basic media new products different from either the enol-keto tautomers or the enolate of TTeA should be generated. In alkaline media, the hydration of the ketone could occur.3

Iglesias

Figure 3. Variation of the absorbance readings of (A) aqueous solutions of 5.5 × 10-4 M TFA [(b) TFA prepared in water; (2) TFA prepared in dioxane] and of (B) aqueous micellar solutions of (b) 1.2 × 10-4 M TFA, [TTACl] ) 0.019 M and (2) 1.35 × 10-4 M TFA, [TTACl] ) 0.040 M as a function of the pH of buffered solutions of acetic acid-acetate, [buffer]t ) 0.076 M. The insets show the corresponding spectra. (‚ ‚ ‚) Simulated curve by assuming E ) 14 000 M-1 cm-1.

Figure 2 displays the spectra of TFA in water ([TFA] ) 1.1 × 10-3 M) and in aqueous micellar solutions of TTACl ([TFA] ) 1.2 × 10-4 M) under several experimental conditions of acid media (HCl or acetic acid-acetate buffer) and of basic media (NaOH or carbonate-bicarbonate buffer). In aqueous hydrochloric acid the spectrum displays a very weak absorption band around 280 nm, which could be due to the enol tautomer of TFA; in the presence of 0.067 M acetic acid-acetate buffer, pH 4.15, the intensity of this absorption band increases; meanwhile, a shift of the maximum toward higher wavelengths is observed, but the absorption intensity is enhanced strongly in the presence of cationic micelles of TTACl (0.044 M), in which case the absorption band maximum (due to enolate) is observed at 295 nm but no effect is observed on the absorption band at 280 nm due to the enol. Note that the concentration of TFA used in the presence of micelles is 10-fold lower than that of pure water. The spectrum recorded at pH 10.20 in 0.052 M carbonate-bicarbonate buffer shows the same trend; a strong absorption band maximum is observed at 295-290 nm either in aqueous micellar solutions or in pure water. Nevertheless, in the presence of micelles, the intensity appears to be higher. We say appears because under these experimental conditions the ketone is unstable, the band at 292 nm decreases with time, but the reaction rate is much higher in the absence of micelles (vide infra). Influence of pH. We analyzed the absorption spectra of both TFA and TTeA as a function of pH. For that we used acetic acid-acetate buffer solutions of total [buffer]

Solutions of the Cationic Surfactant TTAX

Figure 4. Absorption spectra of TTeA (6.4 × 10-5 M) obtained (A) in water and (B) in aqueous micellar solutions, [TTACl] ) 0.019 M, as a function of the pH of buffered solutions of 0.075 M acetic acid-acetate. The insets show the absorbance readings at 341 nm as a function of pH.

) 0.076 M. The spectral study has been performed in pure water and in aqueous micellar solutions of TTACl. Typical results are despicted in Figures 3 and 4 for the cases of TFA and TTeA, respectively. Figure 3 shows two sets of experiments; in one the TFA solution was prepared in water, while in the other TFA was added to the sample mixture from a dioxane solution (41 µL in the absence of TTACl and 10 µL in the presence of TTACl). Notice that the [TFA] used in water is more than 4-fold higher than that used in the presence of the cationic micelles. For both trifluoro-diketones it is observed that increasing the pH also increases the intensity of the absorption band due to the enolate (maximum at 290 or 295 nm for TFA or at 341 nm for the case of TTeA). In the presence of TTACl, the λmax shifts to the red and the absorption intensity is strongly enhanced with respect to the value obtained in water at the same pH. These findings indicate the formation of enolate ions by the presence of cationic micelles. Apparently the electrostatic attractions between enolate anions and the positively charged micellar surface displace the acid ionization equilibrium of keto-enol toward the enolate anions, due to the stabilization of the latter. The clear shift in λmax by the presence of TTACl indicates that the enolate anions in aqueous micellar solutions reside in a microenvironment different from that of bulk water, that is, in the micellar interface, whose dielectric constant in sensibly lower ( ∼ 29-56)15 than that of water ( ) 78.3).16 On the other hand, the well(15) Drummond, C. J.; Grieser, F.; Healy, T. W. J. Phys. Chem. 1988, 92, 2604. (16) Reichardt, C. Solvents and Solvent Effects in Organic Chemistry, 2nd ed.; VCH Verlagsgesellschaft: Weinheim, 1988.

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Figure 5. Absorbance readings at 295 nm of aqueous micellar solutions of TFA (1.03 × 10-4 M), 0.067 M acetic acid-acetate buffer of pH 4.63, measured as a function of surfactant concentration. The insets show the corresponding absorption spectra.

defined isosbestic points at 253 and 308 nm in the spectrum corresponding to TTeA indicate an equilibrium process between at least two species, the keto form and the enolate anion. Influence of Surfactant Concentration. At a constant pH value (acetic acid-acetate buffer) we analyzed the effect of the cationic surfactants TTACl and TTABr on the absorption spectra of both TFA and TTeA. Figures 5 and 6 show typical experimental data. At surfactant concentrations above the critical micelle concentration, cmc, the absorption band due to enolate anion increases strongly with the surfactant concentration. It should be stated that, in the case of TFA, the saturation level is obtained at the total surfactant concentration ∼0.2 M, whereas, with TTeA, concentrations of the surfactant 10fold lower are enough to reach the saturation level. Furthermore, a higher increase in surfactant concentration causes a decrease of the absorbance. These observations are consistent with the enolate association with a micellar interphase due to both electrostatic and specific interactions, the latter being more important with the most hydrophobic substrate, TTeA. In addition, as the species adsorbed to the micellar interphase is an anion, it is possible to observe an exchange process with the counterions of the surfactant when the concentration of the latter was important. If this is so, the displacement of enolate anions from the micellar interphase to the bulk aqueous phase should decrease the absorbance, because in the bulk aqueous phase the enolate becomes protonated. Influence of Counterion Concentration. At fixed pH and fixed surfactant concentration, we analyzed the

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Figure 6. Absorbance readings of aqueous micellar solutions of TTeA (6.4 × 10-5 M), 0.067 M acetic acid-acetate buffer of pH 4.06, measured as a function of surfactant cocentration. The insets show the corresponding absorption spectra.

effect of [Cl-], in the case of TTACl, or [Br-], in the case of TTABr. Typical results are plotted in Figures 7 and 8, showing the effect of [Cl-] and [Br-] on absorbance at λmax of TFA and TTeA spectra, respectively, recorded at fixed surfactant concentration. It is seen that the increase in [Cl-] or [Br-] is accompanied by a decrease in the maximum absorption intensity. The absorbance readings were taken at the maximum wavelength absorption.

Iglesias

Figure 7. Absorbance readings at 295 nm of aqueous micellar solutions of TFA: (A) [TFA] ) 1.03 × 10-4 M, [TTACl] ) 0.050 M, 0.067 M acetic acid-acetate buffer of pH 4.60, measured as a function of [Cl-]; (B) [TFA] ) 1.35 × 10-4 M, [TTABr] ) 0.050 M, 0.067 M acetic acid-acetate buffer of pH 4.60, measured as a function of [Br-]. Scheme 1

Discussion The UV-vis absorption spectrum of 1-phenyl-1,3butadione (benzoylacetone, BZA) in water shows a strong absorption band at 250 and a less intense broad band centered at 312 nm. In cyclohexane solvent (CyH), when BZA is completely in the enol form, the strong absorption band maximum is observed at 306 nm, while the weak absorption band maximum is observed at 245 nm. However, in aqueous alkaline medium (enolate formation) the strong absorption band maximum is observed at 320 nm.12 Both absorption bands observed in CyH solvent must be due to the enol form, for which several resonant forms can be written if the charge-separation structures are considered (see Scheme 1). With acetylacetone (pentane-2,4-dione) there is no possibility of charge-separation structures, and only one absorption band due to enol appears at 274 nm (EH ) 6100 M-1 cm-1 in water and 11 000 M-1 cm-1 in CyH solvent).13 In alkaline medium the absorption band maximum, due to enolate anion, is observed at 292 nm (E ) 23 000 M-1 cm-1). The enolate is perfectly stable,

generated either in alkaline medium or in carbonatebicarbonate buffer, and also in the absence or in the presence of cationic micelles. These preliminar considerations lead to a better understanding of the spectra shown in Figures 1 and 2 for TTeA and TFA, respectively. The spectrum of TTeA resembles that of BZA if one looks at the number of absorption bands and their relative positions. The structure of TFA resembles that of AcAc, and consequently the spectra are similar except in the intensity of the absorption band due to enol, because of the minor proportion of the enol of TFA in water (vide infra). In the presence of

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species in the sense that there are two sites of protonation and the equilibrium enolization constant, KE, linking the two tautomeric protonated forms is a function of the acidity constants of these two species. Thus, titration of an equilibrium mixture of keto (KH) and enol (EH) tautomers yields a mixed acid ionization constant whose value corresponds to the pH of the inflection point (Scheme 2). The acid ionization constant measured in water can be defined as

Ka )

[E-]t[H+] [KH] + [EH]

)

KaKKaE KaK + KaE

) KaEKE KaK ) (1) 1 + KE 1 + KE

Obviously, if KE , 1, then Ka ≈ KaK and KaE . KaK; that is, the minor component of a pair of compounds in equilibrium with a common anion is the stronger acid. From eq 1, taking into account that Aλ ) EHl[EH] + El[E-] (with λ ) 290 or 341 nm for the case of TFA or TTeA in water, respectively) and that [ketone]t ) [keto] + [enol] + [enolate], one can arrive at eq 2 , which relates the

A290,341 )

Figure 8. Absorbance readings at 341 nm of aqueous micellar solutions of 6.4 × 10-5 M TTeA: (A) [TTACl] ) 0.022 M, 0.075 M acetic acid-acetate buffer of pH 4.06, measured as a function of [Cl-]; (B) [TTABr] ) 0.018 M, 0.067 M acetic acid-acetate buffer of pH 4.06, measured as a function of [Br-]. Scheme 2

10pKa + 10pH

(2)

absorbance at the maximum wavelength to the pH of the aqueous solution. In this equation Aa and Ab refer to the absorbance readings at very low and at very high pH, respectively. The solid sigmoid curves in Figures 3 and 4 correspond to the fit of eq 2 to the experimental data using the parameters reported in Table 1, and with Ab and pKa being the adjustable parameters, to avoid problems with the decomposition of enolate at high pH. Once Ab ()A∞E) is known, it is possible to fit the experimental data Aλ - [H+] to eq 3. This procedure allows us to determine the enol acidity constant, that is, the pKaE, and the enolization equilibrium constant KE.

A290 )

carbonate-bicarbonate buffer the absorption of the enolate of TFA at 292 nm is strong, but contrarily to the case of AcAc, the enolate is unstable. The decreasing absorbance at 292 nm with time fits perfectly first-order kinetics (r > 0.9999) in the range of [TFA] ) (0.6-1.4) × 10-4 M. The observed rate constant in 0.065 M carbonate-bicarbonate buffer at pH 10.10 has been measured as k0 ) 7.5 × 10-5 s-1, is independent of diketone concentration, and increases very little with [buffer]. The extrapolated initial absorbance readings gave a value of the extinction coefficient E ) 11 300 M-1 cm-1 for TFA-enolate. The presence of cationic micelles strongly stabilizes the enolate by electrostatic interactions. In addition, the water content at the micellar interface is much lower than that in bulk water. Consequently, the enolate decomposition reaction rate is much less in the presence of TTACl micelles. The dependence of absorbance measurements on pH describes a sigmoid curve. Enolate anions are ambient

Aa 10pKa + Ab 10pH

A∞EH[H+] + A∞EKaE 1 + KE + KaE + [H ] KE

(3)

The solid lines in Figure 9 fit A290 against [H+] using (1 + KE)/KE ) 83 ( 3; KaE ) (2.56 ( 0.18) × 10-4 M; A∞EKaE ) (5.6 ( 0.2) × 10-4 M; and A∞EH ) 3.7 ( 1. These results give KE ) 0.0122 ( 0.0004, which compares quite well with published values;4 pKaE ) 3.59 and pKa ) 5.51 ≈ pKaK; EH ) 6710 M-1 cm-1, in complete agreement with the absorption coefficient of the enol of AcAc at 274 nm,13 and E ) 4000 ( 300 M-1 cm-1. The latter value is smaller than that expected (∼14 000 M-1 cm-1), but if we try to fit the experimental data with the expected value of A∞E ≈ 7.7, the dotted line in Figure 3a is obtained; that is, the fit is not good at all. However, when TFA was added (41 µL) to the sample mixture (3 mL) from a stock dioxane solution (0.0406 M), the obtained results were (1 + KE)/KE ) 89 ( 5; KaE ) (4.25 ( 0.6) × 10-5 M; A∞EKaE ) (3.65 ( 0.08) × 10-4 M; and A∞EH ) 4 ( 1. From these results E ) 15 600 ( 340 M-1 cm-1, pKaE ) 4.37, and pKa ) 6.30 ≈ pKaK; that is, A∞E takes the expected value. When the spectral changes in the absorption spectrum of TFA due to the pH variation are measured in the presence of cationic micelles, it is necessary to define the

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Iglesias

Table 1. Experimental Conditions and Parameters Obtained in the Study of the Influence of pH on the Absorption Spectra of 1,1,1-Trifluoroacetylacetone (TFA) and 1,1,1-Trifluoro-3-(2-thenoyl)acetone (TTeA) in Water and in Aqueous Micellar Solutions of TTACl Experimental Data, Aλ-pH, Fitted to Eq 2 [ketone]/M

medium

10-4

TFA, 5.5 × TFA, 5.5 × 10-4 b TFA, 1.2 × 10-4 TFA, 1.35 × 10-4 b TTeA, 6.4 × 10-5 TTeA, 6.4 × 10-5

watera watera TTAClc TTACld watera TTAClc

λ/nm

Aa

Ab

E/M-1 cm-1

pKa

pKaf

rg

290 292 295 295 341 341

0.052 0.027 0.023 0.027 0.0135 0.0132

2.28 ( 0.03 8.6 ( 1 1.60 ( 0.02 2.57 ( 0.04 1.12 ( 0.02 1.11 ( 0.01

4150 15600 13300 19000 17500 17300

5.51 ( 0.03 6.32 ( 0.07 4.85 ( 0.02e 4.82 ( 0.03e 6.16 ( 0.02 3.33 ( 0.02e

6.30 6.30

0.999 0.9996 0.9995 0.9993 0.999 0.998

6.22

Experimental Data, Aλ-[H+], Fitted to Eq 3 [ketone]/M 10-4

medium

TFA, 5.5 × TFA, 5.5 × 10-4 b TFA, 1.2 × 10-4 TFA, 1.35 × 10-4 b TTeA, 6.4 × 10-5 TTeA, 6.4 × 10-5

λ/nm

watera watera TTAClc TTACld watera TTAClc

290 292 295 295 341 341

AE∞KaE/M 10-4

(5.6 ( 0.2) × (3.65 ( 0.08) × 10-4 (19.5 ( 0.4) × 10-4 (35.6 ( 0.9) × 10-4 (0.47 ( 0.02) × 10-4 0.0327 ( 0.0016

KaE/10-4 M

1 + 1/KE

pKa

KE

rg

2.6 ( 0.425 ( 0.06 11.9 ( 0.2 14.0 ( 0.4 0.625 ( 0.05h 296 ( 3

83 ( 3 89 ( 5 84 ( 5 88.5 ( 7 53.8 ( 0.8 63 ( 3

5.51 6.33 4.84e 4.82e 5.96 3.33e

0.012 0.011 0.012 0.011 0.019 0.016

0.998 0.9997 0.9993 0.9998 0.9995 0.999

0.2h

a In all cases the total buffer concentration was 0.076 M. b Stock dioxane solution of TFA. c Aqueous micellar solutions of TTACl of total f g concentration 0.019 M. d [TTACl])0.040 M. e Refers to the pKap a defined in eq 4. Literature values from ref 4. Correlation coefficient of the goodness-of-fit. h Literature values4 are 2.0 × 10-3 M and 5.0 × 10-5 M for TFA and TTeA, respectively.

Figure 9. Absorbance readings at the maximum wavelength absorption as a function of [H+] of (a) TFA in water, (inset) TFA in the presence of TTACl, (b) TTeA in the presence of TTACl, and (inset) TTeA in water. Solid lines fit eq 3; for parameters see Table 1.

apparent acid ionization constant given in eq 4 , for which

Kaap )

([E-]w + [E-]m)[H+] [KH] + [EH]

)

[E-]t[H+] [KH] + [EH]

(4)

we are neglecting the possible [EH] dissolved in the micellar phase. The influence of the pH in the presence

of a fixed amount of the cationic surfactant TTACl gives the results reported in Table 1. Note that in the presence of TTACl the Ab parameter reaches the expected value, which can be ascribed to the stability of the enolate in the micellar interface. On the other hand, the pKaap is lower than the pKa; in other words, the acidity of TFA increases due to the presence of TTACl micelles. In obtaining eq 3 we assumed that both the enol and enolate forms contribute to the absorption, which might be a necessary statement in the study of the influence of pH in water upon the absorption spectrum of TFA due to the high ketone concentration used. However, in the other situations the fitting process of the experimental data to this equation is insensitive to the value of A∞EH. This fact indicates that the enol concentration present in both the aqueous and micellar pseudophases is negligible. Thus, we assigned A∞EH ) 0. The values so determined for the remaining parameters in eq 3 are reported in Table 1. As we can see, the KE values determined both in water and in aqueous micellar medium are identical; that is to say, the enol form does not associate to micelles. This statement is further corrobotated from the null spectral changes observed in strong mineral acid (no possibility of enolate formation) on varying the cationic surfactant concentration. Proceeding in the same way, the influence of pH on the absorption spectrum of TTeA has been treated as above, and the determined parameters are also given in Table 1. Note that the pKa (or pKaap, in the presence of TTACl), obtained directly from eq 2 or calculated from KE, and the KaE obtained from eq 3 are in agreement, a fact which further validates our previous assumption that the association of the enol form with micelles is negligible. An important observation is that the enol forms of trifluoro derivatives of AcAc are much more acidic than the keto forms (pKa ≈ pKaK). This characteristic is a consequence of the stabilization of a negative charge on the carbonyl group adjacent to the CF3 substituent, due to the high electronegativity of F atoms. As an illustrative result, we might remember that while CF3CH2OH has pKa ) 12.4 in water,17 ethanol is not ionized in water. Finally, in disagreement with various statements found in the (17) Jencks, W. P. J. Am. Chem. Soc. 1979, 101, 5774.

Solutions of the Cationic Surfactant TTAX

Langmuir, Vol. 16, No. 22, 2000 8445

Table 2. Experimental Conditions and Parameters Obtained in the Study of the Influence of Surfactant Concentration on the Absorption Spectrum of TFA (1.03 × 10-4 M, Dissolved in an Aqueous Solution of Acetic Acid-Acetate Buffer of Total Concentration 0.068 M, pH 4.60) and of TTeA (6.4 × 10-5 M Dissolved in an Aqueous Solution of 0.068 M Acetic Acid-Acetate Buffer, pH 4.08) with Experimental Data, Aλ versus [TTAX], Fitted to Eq 5 surfactant

ketone

103cmc/M

AE∞

E/dm3 mol-1 cm-1

TTACl TTABr TTACl TTABr

TFA TFA TTeA TTeA

3.2 2.6 2.7 2.5

0.850 ( 0.004 0.552 ( 0.002 1.12 ( 0.03 1.15 ( 0.04

8250a 5360a 17500 17970

a

1 + [H+]/Ka 12.1 ( 0.8 10.1 ( 0.7 112 ( 11 125 ( 12

Ks/dm3 mol-1

pKa

r

796 ( 64 877 ( 78 (4.2 ( 0.4) × 104 (5.8 ( 0.8) × 104

5.64 5.56 6.12 6.17

0.9997 0.9996 0.997 0.995

Unexpected values due to the enolate anion exchange with surfactant counterions and ketone decomposition in water.

and [X-]ad being the concentration of Cl- of Br- ions added to the medium as the sodium salts. From the definition of Ka, that [ketone]t ) [KH] + [EH] + [E-]w + [E-]m, and that Aλ ) El[E-]t, it is not difficult to obtain eq 7 , which relates the absorbance measure-

Scheme 3

Aλ ) KaAmax λ +

([X-]ad + [TTAX]t + (KI - 1)β[TTAX]m)

Ka + [H ] literature,2-5,10 1,1,1-trifluoroacetylacetone and 1,1,1trifluoro-3-(2-thenoyl)acetone are less than 2% enolized in water. The fluorination degree would increase the enol content of AcAc derivatives solely when H-bonding formation between the OH group of the enol and the F atoms is possible. The influence of surfactant concentration at constant pH was quantitatively explained on the basis of Scheme 3. We have considered that the diketone exists quantitatively as the keto tautomer, neither the keto nor the enol tautomers associate with micelles, and the enhancement of the absorbance by surfactant addition is due to the enolate concentration increase, which we assume to be dissolved in the highly hydrated Stern layer, and that the absorption coefficients in both regions are equal. Under these considerations Aλ ) El([E-]w + [E-]m), where the subscritps w and m mean water and micellar pseudophases, respectively. Then, taking into account Scheme 3, one easily arrives at eq 5 , which relates the absorbance

Aλ )

A∞E(1 + Ks[TTAX]m) Ka + [H+] + Ks[TTAX]m Ka

(5)

to surfactant concentration. The solid lines in Figures 5 and 6 correspond to the fit of the experimental A-[TTAX] data to this equation by using the parameters reported in Table 2.The calculated pKa values agree with those obtained from the A-pH analysis. We note the higher values of Ks determined for the enolate of TTeA (∼5 × 104 mol-1 dm3) than for the enolate of TFA (∼800 mol-1 dm3). This result reflects the higher hydrophobic character of TTeA in comparison with TFA. The exchange process between enolate anions (E-) in the micellar phase and the surfactant counterions Cl- or Br- can be quantified through the equilibrium in eq 6. X Ew + Xm h Em + Xw; KI ≡ KE

(6)

Since the concentrations of TFA ()1.03 × 10-4 M) and TTeA ()6.4 × 10-5 M) used in the study of counterion effects are much lower than [X-], we can approximate [X-]m ) β[TTAX]m and [X-]w ) [X-]ad + cmc + (1 - β)[TTAX]m, with β being the micellar charge neutralized18 (18) Iglesias, E. Langmuir 1998, 14, 5764.

[X-]ad + [TTAX]t +

(

KIKa

Ka + [H+]

)

)

- 1 β[TTAX]m a + b[X-]ad 1 + c[X-]ad

(7)

ments to [X-]ad. Experimental values of Aλ versus [X-]ad have been fitted to this equation. The solid lines in Figure 7 are calculated points from eq 7 using the values of the fitting parameters a, b, and c reported in Table 3. From values of c-1 ) [TTAX]t + β[TTAX]m{1 - KIKa/(Ka + [H+])}, or alternatively from the ratio of a/b ) [TTAX]t + (KI 1)β[TTAX]m, it is possible to determine KI, the exchange equilibrium constant between enolate ions in water and the counterions in the micellar phase. For the equilibrium exchange of the enolate of TFA with Cl- we obtain KI ≡ KCl E ) 29 ( 1, whereas, the exchange with Br ions gives Cl Br Cl ) 8.5 ( 0.4. Thus, K / K ≡ K ()3.4) is the KI ≡ KBr E E E Br + Cl h Br equilibrium exchange constant for Brw m m + 19 Clw, whose value is reported in the literature as 5. The experiments with TTeA were done at [TTAX] ) 0.022 M (corresponding to the maximum absorbance increase) and pH 4.1; that is, Ka , [H+]. The enolate concentration in water is negligible (see inset of Figure 4); that is, A341 ) El[E-]m. These considerations result in an equation like (7), in which c-1 ) [TTAX]t + β[TTAX]m{(KIKa/[H+]) - 1}; b ) 0; and a/c ) (Ka/[H+]) Amax 341 KIβ[TTAX]m. From the values of a and c listed in Table 3, we determine KI ≡ KCl E ) 1750 ( 130 for the equilibrium exchange between Cl- ions and the enolate ions of TTeA, and KI ≡ KBr E ) 450 ( 50 for the equilibrium exchange with Br-. These values give KCBrl ) 3.9, which compares well with the reported value of 5.19 At this point, we can understand also the change in pKa (∆pKa ) pKa - pKaap) caused by the presence of cationic micelles. Starting with eq 4 and taking into account the definition of Ks (Scheme 3), we get Kaap ) Ka(1 + Ks[TTACl]m). For TTeA we determined pKaap ) 3.33 at [TTACl] ) 0.019 M, corresponding to the maximum change in pKa (i.e. to the maximum absorbance variation). Then, we can also calculate Ks ) 4.2 × 104 mol-1 dm3 by using the known parameters occurring in the above expression reported in Table 1. This result matches the value (19) Bartet, D.; Gamboa, C.; Sepu´lveda, L. J. Phys. Chem. 1980, 84, 272.

8446

Langmuir, Vol. 16, No. 22, 2000

Iglesias

Table 3. Experimental Conditions and Parameters Obtained in the Study of the Influence of Cl- or Br- Concentration on the Absorption Spectra of TFA (1.03 × 10-4 M) and of TTeA (6.4 × 10-5 M) Recorded at Fixed Amounts of Surfactant in Aqueus Buffered Solutions of Acetic Acid-Acetate ketone

[TTAX]/M

a

b/dm3 mol-1

c/dm3 mol-1

r

β18

103cmc/M

KId

TFAa TFAa TTeAb TTeAb

Cl;c 0.050 Br;c 0.050 Cl;c 0.022 Br;c 0.022

0.674 ( 0.005 0.562 ( 0.003 0.885 ( 0.010 0.874 ( 0.005

0.76 ( 0.03 1.25 ( 0.07

10.2 ( 0.4 21.8 ( 0.4 3.8 ( 0.2 11.6 ( 0.3

0.9992 0.9994 0.9998 0.9996

0.66 0.76 0.66 0.76

3.2 2.0 3.2 2.0

29 8.5 1750 450

a Acetic acid-acetate buffer, pH 4.6; [buffer] ) 0.068 M. b Acetic acid-acetate buffer, pH 4.06; [buffer] ) 0.076 M. c X in TTAX. t t value of those calcuted from c-1 or a/b.

determined in the study of [surfactant] variation at fixed pH. In the case of TFA, pKaap ) 4.82 determined at [TTACl] ) 0.040 M, together with Ka ) 4.62 × 10-7 mol dm-3 determined from the influence of pH upon A292 when TFA was added from the dioxane solution, gives Ks ) 860 mol-1 dm3, in perfect agreement with the Ks value determined from the influence of surfactant. From Ka and Kaap values one determines ∆pKa ) 2.82 with TTeA while ∆pKa ) 1.50 with TFA. The effect is stronger in the case of TTeA because the association of enolate anions to the positively charged micellar surface is not only due to electrostatic interactions but also due to specific interactions, the latter being higher for the enolate ions of TTeA as a consequence of their higher hydrophobic character. The electrostatic contribution may be deduced from the maximum variation of pKa,20,21 that is ∆µeo ) FΦ ) 2.303RT(pKa - pKaap), where F is the Faraday constant and Φ is the surface potential of the micelle. The above values give Φ ) 156 ( 3 and 93 ( 3 mV for the electrical potential sensed in the micelle by the enolate anion of TTeA and TFA, respectively. These different values could be a consequence of the different locations of both anions in the micellar interphase. In the lyotropic series of anions,22 their different locations in the micellar interphase are often explained as caused by their different hydration levels and sizes. Highly hydrated ions are weakly bound and therefore may remain at a distance where the surface potential would be less than that attributed to the micelle surface (e.g. +148 mV for CTABr of TTABr),23 or they may enter the Stern layer, losing water of hydration. In this way the nonhydrated ions would be bound exclusively by electrostatic forces. Considering the ionic exchange as a transfer free energy from water to micelles for the anions involved in the exchange for Cl- or Br-, X-, and enolate anions, E-, of TFA and TTeA, we may use eq 8 to relate the equilibrium (20) Ferna´ndez, A.; Iglesias, E.; Garcı´a-Rı´o, L.; Leis, J. R. Langmuir 1995, 11, 1917. (21) Gamboa, C.; Sepu´lveda, L.; Soto, R. J. Phys. Chem. 1981, 85, 1429. (22) Larsen, J. W.; Magid, L. J. J. Am. Chem. Soc. 1974, 96, 5774. (23) Ferna´ndez, M. S.; Fromherz, P. J. Phys. Chem. 1977, 81, 1755.

d

Mean

exchange constant (KI) with the total free energy of transfer, ∆µ°enolate.

∆µ°enolate ) ∆µ°X- + RT ln KI ) (∆µ°E)e + (∆µ°E)s (8) As ∆µ°X refers to the free energy of transfer of X- between water and micelles, it takes the same value for TFA enolate as for TTeA enolate. Then from KI ()KCl E ) corresponding to the exchange process between enolate ions of TFA or of TTeA, respectively, and Cl- ions, along with the electrostatic contributions (∆µ°E)e calculated in each case, it is possible to obtain (∆µ°TTeA)s - (∆µ°TFA)s ) -7.90 kJ/mol for the difference in the free energy of transfer by hydrophobic effects between both enolate ions. The negative value accounts for the higher hydrophobicity of TTeA enolate compared to the enolate of TFA. Conclusions In the present investigation it has been shown that the acid ionization equilibrium of 1,1,1-trifluoroacetylacetone and 1,1,1-trifluoro-3-(3-thenoyl)acetone is strongly modified by the presence of cationic micelles. Enolate ions bind to cationic micelles, and the acidity of the corresponding trifluoro-diketones increases. The quantitative analysis of the spectral changes caused by pH variation allows the determination of both the enolization equilibrium constant and the acid ionization equilibrium constant. The results indicate that both diketones are less than 2% enolized in water, that the presence of micelles does not increase the enol fraction, and that the enol form is more acidic than the keto form. The presence of cationic micelles enhances the acidic character of both ketones. At high counterions concentration, the ion exchange process between enolate ions and the counterions is observed. The exchange results not only from electrostatic interactions but also from hydrophobic interactions. Acknowledgment. Financial support from the Direccio´n General de Investigacio´n Cientı´fica y Te´cnica of Spain (Ministry of Education and Science), proyect PB961085, is gratefully acknowledged. LA000800N