The Hammett acidity function H0 in trifluoroacetic acid–dichloromethane mixtures

Tetrahedron Letters 55 (2014) 4325–4327

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The Hammett acidity function H0 in trifluoroacetic acid–dichloromethane mixtures Elena E. Suslova ⇑, Ekaterina N. Ovchenkova, Tatyana N. Lomova G.A. Krestov Institute of Solution Chemistry of the Russian Academy of Sciences, 1 Academicheskaya Str., 153045 Ivanovo, Russian Federation

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Article history: Received 1 April 2014 Revised 27 May 2014 Accepted 5 June 2014 Available online 12 June 2014 Keywords: Trifluoroacetic acid Dichloromethane mixtures Acidity function 2,9,16,23-tert-Butylphthalocyanine Basicity

a b s t r a c t The acidity function (H0) of solutions of trifluoroacetic acid (TFA) in dichloromethane was measured by the indicator method at 298 K in the whole concentration range. The H0 value for the most acidic solution studied (12.93 M trifluoroacetic acid) is 3.09. The equation describing the dependence of H0 on the acid concentration was determined. The obtained quantitative data were used for a spectrophotometric study of the basicity of 2,9,16,23-tetra-tert-butylphthalocyanine. Two forms with different UV/vis spectra were observed and their stability constants determined. Ó 2014 Elsevier Ltd. All rights reserved.

Solutions of trifluoroacetic acid in dichloromethane are of significant practical and theoretical interest. These solutions are reported to remove amino group protection in peptide synthesis1 and are often used for studying macroheterocyclic compound basicity.2,3 The acidity function has been determined for strong acid–water systems,4 including TFA–H2O.5 Acid–base interactions in aqueous solutions of strong acids (HA) lead to ions, the composition, and structure of which have been well-studied.6 In nonaqueous solutions of acids, the HA–solvent interaction is often limited to the formation of ion pairs with partial transfer of a proton7,8 or of molecular complexes,9 that greatly complicates solution quantitative composition determination. Standard indicator ionization mechanisms are used to determine acidity function change during the transition from water to aprotic solvents.10 In addition, the difference in solvation of weak organic bases (indicators and reagents) in water and organic solvents changes their basicities. Chemistry technology developments and science problems require that reliable quantitative data on the compositions of mixed protolytic organic solvents be obtained. The acid–base properties of porphyrazines of various structures were recently studied in a mixed TFA–CH2Cl2 solvent.11 However, the authors identified the concentration stability constants of the acid forms only, since the values of the acidity function of this medium remain unknown. This is why correlation analysis of ‘structure–property’ using data for different solvents is complicated. ⇑ Corresponding author. Tel.: +7 4932336990; fax: +7 4932336237. E-mail addresses: [email protected], [email protected] (E.E. Suslova). http://dx.doi.org/10.1016/j.tetlet.2014.06.021 0040-4039/Ó 2014 Elsevier Ltd. All rights reserved.

The purpose of this work was to obtain the acidity function of trifluoroacetic acid solutions in dichloromethane over a range of mixture compositions, and to study the acid–base equilibrium of 2,9,16,23-tetra-tert-butylphthalocyanine [H2Pc(t-Bu)4] (Fig. 1) in this medium. The concentrations of the non-ionized forms of indicators I–IV (see Table 1) in solutions at various concentrations of TFA in dichloromethane were determined spectrophotometrically at 298 K. The UV/vis spectra were recorded on an Agilent 8453 spectrophotometer. Measurements were performed at the absorption maxima of the non-ionized forms of I–IV, in the visible spectral range. The numerical values of the molar extinction coefficients, eN and eD, which were necessary for concentration ratio (indicator ratio I) calculations of the non-ionized (N) and ionized (D) forms of I–IV were obtained as follows. The values of eN were determined from the optical densities of solutions without the acid, and those

t-Bu t-Bu

N NH

N

N

N N

HN N

t-Bu

t-Bu

Figure 1. Structure of 2,9,16,23-tetra-tert-butylphthalocyanine.

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Table 1 Stability constants of indicators (pK) and concentration intervals of TFA in dichloromethane Indicator

CTFA, M

pK

3-Nitroaniline (I) 4-Nitroaniline (II) 2-Nitroaniline (III) 4-Nitrodiphenylamine (IV)

0.006–0.13 0.034–0.21 0.14–0.71 0.23–1.70

0.43 0.59 ± 0.05 1.44 ± 0.07 2.03 ± 0.19

invariable and equal to 0.43 (Table 1) over the concentration range of 0.006–0.13 mol/L TFA. The pK values for II–IV are also presented in Table 1. The acidity scale of TFA solutions in CH2Cl2 at 298 K obtained using the I–IV pK and lgI values is presented in Figure 2. The dependence of the acidity function H0 on TFA concentrations is described by Eq. 3 with a correlation coefficient R2 = 0.998.

H0 ¼ 2:3705  ex=0:0729 þ 1:4866  ex=0:5177  0:2932  ex=7:2917  1:4819 of eD were determined in solutions of TFA at concentrations that provided complete ionization of these indicators. The acidity function (H0) of the TFA–CH2Cl2 system was calculated from the Hammett Eq. 1.

H0 ¼ pK þ lgI

ð1Þ

The concentration ranges of the acid at which the values of lgI and pK for I–IV were obtained are given in Table 1. Acid–base interactions of H2Pc(t-Bu)4 were studied by spectrophotometric titrations at 298 K, following the described procedure.12 The concentration ratio of equilibrated acid and base forms was calculated using Eq. 2.

I ¼ ðAi  A0 Þ=ðA1  Ai Þ

ð2Þ

Here, A0 is the initial optical density of a solution at an analytical wavelength, A1 is the optical density corresponding to the complete transformation of the base into protonated form, Ai is the optical density in i-th run. The number of protons involved in the acid–base interaction (n) was determined by optimization of the linear dependences, lgI–H0. The stability constants of the protonated forms (K1 and K2) were calculated by the least-squares method with use of Microsoft Excel. The relative error in the determination of K1 and K2 does not exceed 25%. The acidity of TFA solutions in dichloromethane was determined using the Hammett method.13 The acidity scale obtained by the indicator method is a thermodynamic characteristic of a solution in standard state that is an infinitely dilute solution of the acid in the solvent studied. The indicator acidity scale may reflect both the protonating ability of the solution14 and some other types of acid–base interaction between the acid and a weak organic base.15 This depends on the indicator ionization mechanism. In order to plot an acidity scale by the indicator method, the stability constant (pK) of the most basic indicators used must be known. The pK values of other indicators were determined by the overlap method. The pK of the solution of indicator I in TFA was equated to the concentration ionization constant, which is

Figure 2. Acidity functions (H0) of TFA–CH2Cl2 (1) and TFA–H2O [5] (2) mixtures at 298 K.

ð3Þ

Here x is the TFA concentration in mol/L. The minimum H0 value is reached at a molar ratio of acid–solvent of approximately 5:1 because of the decrease in the ionizing power of the trifluoroacetic acid molecules on their self-association in the TFA–CH2Cl2 system (Fig. 2, line 1) as in the TFA–H2O system.5 In 100% acid, both TFA dimers and associates of higher orders are formed.16 The destruction of associates occurs on addition of dichloromethane, which leads to an increase in the acidity of the solution and the formation of molecular complexes with CH2Cl2 with lower ionizing properties than those of TFA molecules. The acidity function values of different homogeneous systems with different acid–base properties are often used for correlating the catalytic activity and acidity. Their absolute values depend on the chosen standard state. This is clearly illustrated by the H0 values of the TFA–CH2Cl2 and TFA–H2O systems (Fig. 2 and Ref. 5, respectively). In these systems the acidity functions (H0) are standardized differently: to infinitely dilute solutions of TFA in dichloromethane and in water, respectively. For 100% TFA, two values of the acidity function were obtained: 3.09 and 2.71 from the H0 scales in CH2Cl2 (Fig. 2, line 1) and in H2O (Fig. 2, line 2). The literature value of H0 is 3.03 for the TFA–H2SO4 system,17 where 100% TFA is accepted as a standard state. Quantitative comparison of the acid–base system catalytic properties necessitated that the acidity functions and basicity constants used for this purpose were standardized in the same way. The acid–base reaction of 2,9,16,23-tetra-tert-butylphthalocyanine in proton-donor solvents was studied as an example of the use of the acidity function equation in TFA–CH2Cl2 systems. H2Pc(t-Bu)4 exists in TFA–CH2Cl2 in several acid–base forms differing in the position and intensity of the absorption maxima in the UV/vis spectrum. This allows individual forms to be observed very easily.

Figure 3. Variation of the electronic absorption spectrum of H2Pc(t-Bu)4 in TFA– CH2Cl2 (H0 1.80–0.29).

E. E. Suslova et al. / Tetrahedron Letters 55 (2014) 4325–4327

Figure 4. Variation of the UV/vis spectra of H2Pc(t-Bu)4 in TFA–CH2Cl2 (H0 0.29 to 2.33).

The UV/vis spectrum of H2Pc(t-Bu)4 in dichloromethane contains two maxima at k 698 (Qx) and 663 (Qy) nm. Addition of trifluoroacetic acid reduces the intensity of these bands and increases the intensity of the absorption maximum at k 742 nm (Fig. 3). The spectrum changes in the range of TFA concentrations from 1.6  103 to 0.08 M. Further increase in the TFA concentration (acidity range H0 from 0.29 to 2.33, CTFA = 0.08–7.76 M, Fig. 4) results in a new absorption maximum at 763 nm and reduces the intensity of the absorption band at 742 nm. The fact that two spectral curve series are obtained indicates that there are two stages for the reaction course. It is known that acid–base interactions of phthalocyanines may involve both meso-nitrogen atoms and internal (pyrrole) nitrogen atoms.18 In our case, a bathochromic shift of the Q-band both in the first range of TFA concentrations and in the second, suggests protonation of the meso-nitrogen atoms. The number of donor centers involved in the acid–base interaction was determined using the straight-line plot of lgI versus H0; its slope is equal to 0.89 (1) and 1.28 (1) for the first and second steps of protonation, respectively. Based on the experimental data, we can derive the Eqs. 4 and 5 for the protonation reaction. K2

H2 Pcðt-BuÞ4 þ Hþ ½H2 Pcðt-BuÞ4 Hþ K2

½H2 Pcðt-BuÞ4 Hþ þ Hþ ½H2 Pcðt-BuÞ4 H2 2þ

ð4Þ ð5Þ

The thermodynamic stability constants of the protonated forms [H2Pc(t-Bu)4H]+ K1 = (0.04 ± 0.01)  102 and [H2Pc(t-Bu)4H2]2+ K2 = (25.2 ± 6.3) mol/L are determined using Eqs. 6 and 7.

pK 1 ¼ nH0 þ lgðC ½H2 PcðtBuÞ4 Hþ =C H2 PcðtBuÞ4 Þ

ð6Þ

pK 2 ¼ nH0 þ lgðC ½H2 PcðtBuÞ

ð7Þ

4 H2 

2þ

=C ½H2 PcðtBuÞ4 Hþ Þ

Thus, the formation of only the first and the second protonated forms of the substituted phthalocyanine is observed in TFA–CH2Cl2. Each protonation stage is attended by a strong feedback in the

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visible range of the UV/vis spectrum that gives the opportunity to ‘operate’ the physicochemical properties of a chromophore by changing the solvent acidity. In fact, the third and fourth protonated forms of H2Pc(t-Bu)4 are observed only in H2SO4–CH3COOH at H0 5.70 and 8.00, respectively.18 The reliability of the above data is confirmed by comparison with those for H2SO4–CH3COOH mixtures. The first and second protonated forms of H2Pc(t-Bu)4 in H2SO4–AcOH mixtures are observed at H0 0.30 and 2.40, respectively.18 In TFA–CH2Cl2, the acidity function H0 for H2Pc(t-Bu)4H+ is 0.29 and for H2Pc (t-Bu)4H2+ 2 is 2.33. Thus, the acidity range for the formation of the first and second protonated forms is the same. In conclusion, we have defined the acidity scale for solutions of TFA in CH2Cl2 over a wide range of concentrations of TFA at 298 K. The obtained data confirm previous results from another solvent system,18 and indicate that our data are correct and accurate. The study of the H2Pc(t-Bu)4 acid–base equilibrium in this medium results in the thermodynamic stability constants of the protonated forms [H2Pc(t-Bu)4H]+ and [H2Pc(t-Bu)4H2]2+ of (0.04 ± 0.01)102 and (25.2 ± 6.3) mol/L, respectively. Acknowledgments Financial support from Grants of the N. 8 Program of Fundamental Research of the Russian Academy of Sciences and the Russian Foundation for Basic Research, Projects No. 12-03-00967 is gratefully acknowledged. References and notes 1. Jakubke, H. D.; Jeschkeit, H. Aminosäuren, Peptide, Proteine; Academie: Berlin, 1982. 2. Saito, S.; Shin, J. Y.; Lim, J. M.; Kim, K. S.; Kim, D.; Osuka, A. Angew. Chem. 2008, 120, 9803–9806. 3. Khelevina, O. G.; Bubnova, A. S.; Makarova, O. N.; Lukina, S. A.; Vagin, S. I.; Stuzhin, P. A. Koord. Khim. 2006, 32, 468–474. Russ. J. Coord. Chem. 2006, 32, 451–457. 4. Paul, M. A.; Long, F. A. Chem. Rev. 1957, 57, 1–45. 5. Spitzer, U. A.; Toone, T. W.; Stewart, R. Can. J. Chem. 1976, 54, 440–447. 6. Yukhnevich, G. V.; Tarakanova, E. G.; Mayorov, V. D.; Librovich, N. B. Uspekhi Khim. 1995, 64, 963–974. Russ. Chem. Rev, 1995, 64, 963–974. 7. Burdin, V. V.; Kislina, I. S.; Maiorov, V. D.; Sysoeva, S. G.; Librovich, N. B. Izv. Akad. Nauk, Ser. Khim. 1998, 47, 2484–2489. Russ. Chem. Bull., Int. Ed. 1998, 47, 2404–2409. 8. Kislina, I. S.; Sysoeva, S. G.; Librovich, N. B.; Temkin, O. N.; Eremenko, I. L.; Nefedov, S. E. Dokl. Akad. Nauk 1998, 360, 649–651. Dokl. Chem. 1998, 360, 106– 108. 9. Maiorov, V. D.; Voloshenko, G. I.; Kirilova, A. P.; Librovich, N. B. Izv. Akad. Nauk, Ser. Khim. 1999, 48, 312–317. Russ. Chem. Bull., Int. Ed. 1999, 48, 313–318. 10. Rochester, C. H. Acidity Functions; Academic Pres: London, 1970; pp 234–264. Chapter 6. 11. Khelevina, O. G.; Bubnova, A. S.; Makarova, O. N. Zh. Obsch. Khim. 2006, 76, 1569–1574. Russ. J. Gen. Chem., 2006, 76, 1504–1509. 12. Albert, A.; Serjeant, E. Ionization Constants of Acids and Bases; Methuen: London, 1962. 13. Hammett, L. P.; Deyrup, A. J. J. Am. Chem. Soc. 1932, 54, 439. 14. Vinnik, M. I.; Kislina, I. S.; Librovich, N. B. Dokl. Akad. Nauk SSSR 1980, 251, 138. Dokl. Chem. 1980, 251, 138. 15. Vinnik, M. I. Kinet. Catal. 1980, 21, 136. 16. Perelygin, I. S.; Afanas‘ev, A. N. J. Struct. Chem. 1973, 14, 1033. 17. Herbert, H. H.; Ronald, A. G. J. Am. Chem. Soc. 1959, 81, 1847–1849. 18. Stuzhin, P. A.; Khelevina, O. G.; Berezin, B. D. In Phthalocyanines Properties and Applications (Azaporphyrins: Acid–base Properties); Leznoff, C. C., Lever, A. B. P., Eds.; VCH: New York, 1996; Vol. 4, pp 9–78.