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• 1a 26 14 Ref. 10 21 57 Ref. 10 90 35 Ref. 10 38 ± 0.03 9.46 ± 0.08 84 ± 0.03 This work 1b 29 ± 0.03 9.38 ± 0.09 68 ± 0.03 This work 1c 05 ± 0.03 9.75 ± 0.08 80 ± 0.02 This work 1d 45 22 Ref. 10 22 56 Ref. 10 95 35 Ref. 10 18 ± 0.01 9.21 ± 0.06 39 ± 0.03 This work 1e - 32 34 Ref. 8 94 ± 0.02 9.16 ± 0.03 11.33 ± 0.08 20.45 ± 0.06 24.43 ± 0.07 This work 1f - 19 40 Ref. 7 46 46 Ref. 9 28 62 11 Ref. 10 11 73 25 Ref. 10 01 61 30 Ref. 10 47 ± 0.04 8.07 ± 0.01 10.76 ± 0.08 18.83 ± 0.07 23.29 ± 0.07 This work 2a 83 ± 0.03 15 ± 0.08 99 ± 0.03 This work 2b 75 ± 0.03 69 ± 0.07 44 ± 0.03 This work 2c 43 ± 0.03 9.15 ± 0.02 58 ± 0.02 This work 2d 07 ± 0.01 9.28 ± 0.02 34 ± 0.03 This work 2e 35 ± 0.02 9.74 ± 0.03 09 ± 0.06 This work 3a 41 25 Ref. 7 58 13 Ref. 8 37 77 Ref. 6 36 ± 0.04 6.09 ± 0.09 9.11 ± 0.01 19 ± 0.04 19.55 ± 0.03 This work 12 Ref. 5 3b 38 ± 0.04 6.54 ± 0.09 10.12 ± 0.04 16.66 ± 0.04 20.03 ± 0.04 This work mM + lL + hH M m L l H h (3) β mlh = [M m L l H h ] [M ] m [L] l [H] h , (4) where M is Cu(II) ion, L is ligand, and H is proton, and m, l, and h are the respective stoichiometric coefficients. The potentiometric data for the Cu(II) – L 2 systems indicate that there is a significant tendency toward the formation of ML 2 species. Cu(II) ion complexes were formed by releasing 2 of the hydrogen ions from the fully protonated form of the ligands. 11 Compounds 1e and 1f serve as tridentate ligands by the coordinating of imino, phenolic –OH and –SH groups with Cu(II) ion. The others (1a, 1b, 1c, 1d, 2a, 2b, 2c, 2d, 2e, 3a, and 3b) serve as bidentate ligands. Thus, the stability constants of Cu(II) complexes of 1e and 1f are higher than those of the others. Coordination numbers of a central atom can change to 6 or 4 depending on the ligand structures in the complex formation. For example, the coordination number of Cu(II) ion was observed as 4 against all studied ligands except 1e and 1f in this study. Therefore, tetrahedral complexes were obtained. On the other hand, the coordination number of Cu(II) ion was 6 against 1e and 1f, because they are tridentate ligands. The molecular structures of 4- and 6-coordinated complexes of copper with 1a and 1f are given in Figure 3a and 3b. Figure 3. Molecular structures of 4- and 6-coordinated complexes of Cu(II) with 1a and 1f (a) Cu(II)–1a 2 (b) Cu(II)–1f 2 Complexes having 6 coordination numbers form an octahedral structure and they can be formulated as MX 2 , where MX 2 is a structure of bis-complex, M is Cu(II) ion, and X is the ligand 1e (or 1f ). Therefore, the octahedral complexes between 1e and 1f and Cu(II) ion are more stable than the tetrahedral complexes. This situation was supported by the experimental (Table 3) and the semiempirical molecule orbital (SE-MO) PM3 method. 3D structures of complex species and their formation heats (H f ) are calculated by PM3 method. Accordingly, formation heats (H f ) of Cu–1e 2 /1f 2 /2a 2 /3b 2 complexes were determined as 608.78 kcal/mol, 642 kcal/mol, 728.96 kcal/mol, and 776.13 kcal/mol, respectively, and the findings are given in Figure 4a–4d. Figure 4. Comparison of 3D structure of complex species and their formation heats (H f ) (a) Cu(II)–1e 2 complex (ONS type) (H f : 6078 kcal/mol) (b) Cu(II)–1f 2 complex (ONO type) (H f : 642 kcal/mol) (c) Cu(II)–2a 2 complex (NO type) (H f : 796 kcal/mol) (d) Cu(II)–3b 2 complex (NS type) (H f : 713 kcal/mol). As a result, different electron densities on the donor atoms are an important factor affecting the stability of the Cu(II) complexes with the ligands. The various complexes between Cu(II) ion and the Schiff bases were formulated as CuL 2 , CuHL 2 , CuH 2 L 2 , and CuH −1 L 2 (Cu (OH) L 2 ) , depending on pH. The overall stability constants of detectable Cu(II)–L 2 species are given in Table 3.
• Table Overall stability constants in Cu(II)–L 2 binary system (25.0 ± 0.1 ◦ C, I : 0.1 M by NaCl, 0.05 mmol HCl). Ligands m h l logβ Ligands m h l logβ 1 2 07 ± 0.05 1 2 88 ± 0.03 1a 1 1 2 38 ± 0.08 2a 1 1 2 90 ± 0.05 1 –1 2 68 ± 0.07 1 2 43 ± 0.05 1 2 18 ± 0.03 1b 1 1 2 97 ± 0.08 2b 1 1 2 96 ± 0.03 1 –1 2 99 ± 0.09 1 –1 2 99 ± 0.07 1c 1 2 65 ± 0.15 2c 1 2 59 ± 0.08 1 –1 2 56 ± 0.10 1 1 2 94 ± 0.15 1d 1 2 24 ± 0.04 2d 1 2 71 ± 0.05 1 –1 2 67 ± 0.09 1 –1 2 45 ± 0.11 1 2 85 ± 0.09 2e 1 2 29 ± 0.04 1e 1 1 2 64 ± 0.05 1 –1 2 84 ± 0.07 1 2 2 55 ± 0.07 3a 1 2 60 ± 0.06 1 2 51 ± 0.05 1 1 2 91 ± 0.04 1f 1 1 2 33 ± 0.04 1 –1 2 76 ± 0.01 1 2 2 11 ± 0.02 3b 1 2 65 ± 0.06 1 1 2 52 ± 0.08 1 –1 2 93 ± 0.08 Electron pairs on donor atoms play a critical role for complex formation. Mobility of the electron pairs facilitates participation in coordination. However, electron-withdrawing groups on the ligands cause decreasing stability in the complexes because of the limitation of electron mobility. This situation is clearly seen in Table Differences in electronegativity of Br and Cl atoms cause different stability constants in the Cu(II)–1a 2 and Cu(II)–1b 2 complexes. The same case can be said for the –OH and –SH groups. The species distribution curves of the complexes between Cu(II) ion and 1e, 1f, 2a, and 3a ligands are given in Figure 5a–5d. In Figure 5, the species distribution curves of 1e differ from those of 1f because of the different electron density of the –OH and –SH groups. In the Cu(II)–1e 2 system, 3 main complexes (CuL 2, CuHL 2 , and CuH 2 L 2 ) were obtained at between pH 5 and 11. The CuHL 2 species start occurring at pH 5 and reach the maximum at pH 8–9 by 90%; and the CuL 2 species start to form at pH 8 and reach the maximum at pH 11 by 99%. In the Cu(II)–1f 2 system, similarly, CuL 2 and CuH 2 L 2 complex species were observed in the acidic and basic region at 99%, the same as in the Cu(II)–1e 2 system. However, CuHL 2 species reach the maximum at pH 8–9 and approx. 60%. In Cu(II)–2a and Cu(II)–3a systems, the main complexes (CuL 2 and CuHL 2 ) were obtained in neutral and acidic regions. The CuL 2 and CuHL 2 species exist above pH 7 at 90% and 98%, respectively. For both complexes (Cu(II)–2a and Cu(II)–3a), hydrolysis species (CuH −1 L 2 ) were also observed at pH 11 and at 99%. The log β CuL2 values are shown in Figure 6. (a) Cu (II) - 1e 2 system (b) Cu (II) - 1f 2 system (c) Cu (II) - 2a 2 system (d) Cu (II) – 3b 2 system % form pH CuHL 2 % form pH 5 6 7 8 9 10 11 % form pH % form pH Figure 5. The species distribution curves of complexes between Cu(II) ion and 1a, 1e, 2a, and 3a ligands (a) Cu (II)–1e 2 system (b) Cu(II)–1f 2 system (c) Cu(II)–2a 2 system (d) Cu(II)–3b 2 system (25.0 ± 0.1 ◦ C, I : 0.1 M by NaCl, 0.05 mmol HCl). 1a 1b 1c 1d 1e 1f 2a 2b 2c 2d 2e 3a 3b 11 12 13 14 15 16 17 18 19 S ta b il it y c o n st an ts o f C u L 2 c o m p le x es Ligands Figure 6. Changing of the log β CuL2 values for the ligands. Experimental procedure and methods Preparation of the Schiff bases The studied Schiff bases were synthesized by a procedure reported by Perumal et al. 30 To a stirred solution of 2-methoxybenzaldehyde (1.36 g, 10 mmol) in ethanol (10 mL) was added a solution of 2-aminothiophenol (1.87 g, 15 mmol) in ethanol (10 mL). The mixture was refluxed for 5 h. After cooling the reaction mixture, the precipitated substance was filtered and recrystallized in ethanol. The other Schiff bases were prepared by the above-mentioned procedure. The physical data of unknown compounds: ( E) -2-(2-methoxybenzylideneamino)benzenethiol (1c): (yield 91%; mp 95–98 ◦ C); 1 H NMR (400 MHz, CDCl 3 ) δg = 8.64 (d, J = 7.6 Hz, 1H), 8.19 (d, J = 8.0 Hz, 1H), 7.98 (d, J = 8.0, Hz, 1H), 7.55 (t, J = 6 Hz, 1H), 7.48 (t, J = 7.2 Hz, 1H), 7.42 (s, 1H, HC = N), 7.28 (t, J = 8.0 Hz, 1H), 7.19 (t, J = 7.6 Hz. 1H9, 05 (d, J = 8.4 Hz, 1H), 4.04 (s, 3H, -OCH 3 ) , 88 (brs, 1H, -SH). 13 C NMR (100 MHz, CDCl 3 ): gδ = 121, 159.16, 157.29, 136.49, 131.87, 129.57, 125.99, 124.68, 122.85, 121.31, 121.20, 116.42, 111.74, 55.89. IR (Liquid): 3544, 3475, 3413, 3226, 1612, 1563, 1415, 1138, 1041, 884, 863, 747, 605, 482. Elemental Anal. Cald: C, 11; H, 5.39; N, 5.76; S, 13.18. Found: C, 68.91; H, 5.27; N, 5.72; S, 13.28. ( E) -2-((1H-pyrrol-2-yl)methyleneamino)benzenethiol (2b): (yield, 73%; mp 76–78 ◦ C); 1 H NMR (400 MHz, CDCl 3 ) δg = 11.74 (s, -NH), 8.36 (s, 1H, HC = N), 7.44 (d, J = 8.0 Hz, 1H), 7.24–7.21 (d, J = 8.0, Hz, 1H), 7.16–7.11 (m, 3H), 6.80 (bs, 1H), 6.24 (m, 1H), 4.56 (s, 1H, SH). 13 C NMR (100 MHz, CDCl 3 ): δ = 122, 149.74, 130.80, 130.61, 127.62, 126.30, 126.02, 125.28, 118.22, 117.89, 110.54. IR (Liquid): 3554, 3482, 3415, 3235, 1616, 1567, 1413, 1132, 1037, 881, 867, 744, 601, 480. Elemental Anal. Calcud: C, 65.32; H, 4.98; N, 85; S, 15.85. Found: C, 65.28; H, 5.18; N, 13.78; S, 15.98.
• Apparatus and materials Firstly, Schiff bases were dissolved in sufficient ethanol and diluted at a ratio of 1/10. Next, 1 × 10 −3 M stock solution was prepared for each ligand. Ethanol, NaCl, and CuCl 2 were purchased from Merck, potassium hydrogen phthalate (KHP) and borax (Na 2 B 4 O 7 ) from Fluka, and 0.1 M NaOH and 0.1 M HCl as standard from Aldrich. All reagents were of analytical quality and were used without further purification. A solution of metal ion (1 × 10 −3 M) was prepared from CuCl 2 as received and standardized with ethylenediaminetetraacetic acid (EDTA). 40 Next, 0 M NaCl stock solution was prepared from the original bottle. For all solutions, CO 2 free double-distilled deionized water was obtained with an aquaMAX-Ultra water purification system (Young Lin Inst.). Its resistivity was 18.2 M Ω cm −1 Potentiometric measurements All potentiometric pH measurements were carried out on solutions in a 100-mL double-walled glass vessel using the Molspin pH meter with Orion 8102BNUWP ROSS ultra combination pH electrode and the temperature was controlled at 25.0 ± 0.1 ◦ C by circulating water through the double-walled glass vessel, from a constanttemperature bath (DIGITERM 100, SELECTA). The electrode was calibrated according to the instructions in the Molspin Manual. 41 An automatic burette was connected to the Molspin pH-mV-meter. The pH electrode was calibrated with a buffer solution of pH 4.005 (KHP) and pH 9.180 (borax) 42 at 0 ( ±0.1) ◦ C. During the titration, nitrogen (99.9%) was purged through the cell. The Hyperquad 43 computer program was used for the calculation of both dissociation and stability constants.
• The cell was equipped with a magnetic stirrer. Atmospheric CO 2 was excluded from the titration cell with a purging steam of purified N 2 . The system was maintained at an ionic strength of 0.1 M by NaCl as supporting electrolyte. A solution containing about 0.01 mmol of the ligands, and the required amount of 1.0 M NaCl and 0.1 M HCl were put into the titration cell. Finally, doubly distilled deionized water was added to the cell to a total volume of 50 mL and titration was started. The pH data points were collected after each addition of 0.03 mL of the standardized NaOH solution. The second solution contained the same amounts of components plus 0.005 mmol of Cu(II) solution and doubly distilled deionized water was added to the same total volume. The potentiometric studies were carried out at the metal:L molar ratios of 1:2 and each titration was repeated 3 times. Conclusion In this work, the effect of substituents on the Cu(II)–L 2 complexes was discussed. Different electron densities on the donor atoms are an important factor affecting the stability of the Cu(II) complexes with the ligands. The various complexes between Cu(II) ion and Schiff bases were formulated as CuL 2 , CuHL 2 , CuH 2 L 2 , and CuH −1 L 2 (Cu (OH) L 2 ) depending on pH. Stability constants of binary complexes between Cu(II) and Schiff bases were determined in 0.1 M ionic strength (NaCl) and at 25.0 ± 0.1 ◦ C, using the combined glass electrode, potentiometrically. The log β CuL2 values for the ligands are shown in Figure 4. The dissociation constants and overall stability constants were calculated using Hyperquad and the results are given in Tables 1 and 2. The coordination number of Cu(II) ion was 4 against all studied ligands except 1e and 1f in this study. Therefore, tetrahedral complexes were obtained. However, the coordination number of Cu(II) ion was 6 against 1e and 1f because they are tridentate ligands. Cu(II)–1e and –1f complexes are in octahedral structure. As a result, the complexes of ONS- and ONO-type tridentate ligands are more stable than those of NO- and NS-type bidentate ligands. The log β CuL2 values are changed as 1e > 1f > 2b > 2a > 3b > 1a > 2d > 1c > 3a > 2c > 1b > 2e > 1d. Acknowledgment The authors gratefully acknowledge the financial support of this work from the Scientific Research Council of Gaziosmanpa¸sa University. References Canpolat, E.; Kaya, M. Turk. J. Chem. 2005, 29, 409–415. Arora, K.; Sharma, K. P. Synth. React. Inorg. Met.-Org. Chem. 2003, 32, 913–919. Vigato, P. A.; Tamburini, S. Coord. Chem. Rev. 2004, 248, 1717–2128. Katsuki, T. Coord. Chem. Rev. 1995, 140, 189–214. Issa, Y.; Fattah, H.; Omar, M.; Solisman, A. Monatsh. Chem. 1995, 126, 163–171. Geary, W.; Nickless, G.; Polland, F. Anal. Chim. Acta 1962, 27, 71–79. Gurkan, P.; Gunduz, N. J. Indian Chem. Soc. 1997, 74, 713–714. Sengupta, G.; Ghosh, N. J. Indian Chem. Soc. 1988, 65, 712–722. Friedrich, A.; Hefele, H.; Monner, A.; Scholz, F. Electroanalysis 1998, 10, 244–248. Demirelli, H.; Koseoglu, F.; Kavak, N. J. Sol. Chem. 2004, 33, 1467–1472. Kumar, R.; Singh, R. Turk. J. Chem. 2006, 30, 77–87.
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