Potentiometric and antimicrobial studies on the asymmetric Schiff bases and their binuclear Ni(II) and Fe(III) complexes; synthesis and characterization of the complexes

2.1 Materials and reagents

Nickel(II) and iron(III) chlorides were obtained from Aldrich Chemical Company and used without further purification. The solvents used were of spectroscopic grade.

2.2 Synthesis

2.3 Instrumentation

Infrared absorption spectra in the 4000–400 cm⁻¹ range were recorded using KBr discs on a Mattson 1000 FTIR spectrophotometer. Carbon, hydrogen and nitrogen analyses were obtained using GmbH vario MICRO CHNS analyzer. The mass spectra (LC–MS/ESI) were recorded at 70 eV on an Agilent 1100 MSD spectrometer. Melting points were measured on a Gallenkamp apparatus using a capillary tube. The TGA curves of complexes were recorded using a Du Pont Instrument 951 between 35 and 800 °C at heating rate of 10 °C min⁻¹ in a N₂ atmosphere. The room temperature magnetic susceptibilities were determined using a Sherwood Scientific MKI model Evans magnetic balance. Conductivity measurements were carried out at 25 °C in 10⁻³ M dimethyl sulfoxide (DMSO) or 10⁻³ M N,N-dimethyl formamide (DMF) using a Siemens WPA CM 35 apparatus. Metal ion solutions were standardized by VARIAN AA 240FS model Fast Sequential Atomic Absorption Spectrometer. Ultraviolet–visible spectra of the complexes in DMSO were recorded using Analytik Jena UV-200 spectrophotometer, over the wavelength range 400–900 nm at room temperature. MIC values of diimine H₂L¹ were determined in Gazi University–Molecular Biology Research and Training Center (MOBAM). XRD was recorded on a APD 2000 PRO X-ray diffractometer with Cu Kα radiation (λ = 1.54059 Å) in the 2θ = 0–70° range.

2.4 Potentiometric reagents and solutions

The potentiometric measurements in this study were carried out in 1: 1 (v: v) aqueous-ethanol media. Ethanol was purchased from Merck and used without further purification. Doubly distilled and CO₂ free water was used exclusively.

Stock solutions of strong acid and strong base were prepared using analytical reagent grade 0.10 M HCl solution (Merck) and 0.10 M KOH solution (Merck), respectively. The 0.0051 M KOH solution was standardized by titration against the primary standard oxalic acid dihydrate (Aldrich). The 0.0506 M HCl solution was prepared and then standardized by titration with 0.0051 M KOH solution. During each titration the ionic strength was maintained at 0.50 M KCl (Merck, extra pure).

All metal ion solutions were prepared from their analytical grade chlorides and standardized by atomic absorption spectroscopy method.

2.5 Potentiometric measurements

All potentiometric measurements were performed in a double walled glass cell that was thermostated at 25.0 ± 0.1 °C. Electrode system was kept in 1:1 (v:v) aqueous-ethanol media. A Jenway pH-meter equipped with an ISE combined glass electrode was used for titrations.

The system was calibrated to read the hydrogen ion concentration by titration of HCl solution with KOH solution in 0.50 M KCl ionic strength according to Gran’s method (Gran, 1952; Sarı et al., 2003). Small amounts of titrant (0.05 mL) KOH solution were added with a microburette. Solutions were titrated between the pH range 3–10 for the ligands, 3–9 for the nickel(II) complexes, but 3–6 for the iron(III) complexes to avoid the precipitations. Concentrations of the nitro-Schiff bases and the asymmetric diimines were between 3.9 × 10⁻⁴ and 4.6 × 10⁻⁴ M and metal-to-ligand ratios were 1:4 for Ni(II) and 1:8 for Fe(III). pH corrections for aqueous-organic system were done as described in the literature (van Uitert and Haas, 1953; Belaid et al., 2007).

The following solutions (total volume = 25.0 ± 0.01 mL) were titrated potentiometrically with 0.0051 M KOH solution:

1. HCl + KCl.

2. HCl + KCl + ligand.

3. HCl + KCl + ligand + metal ion.

The computations of the protonation constants of the nitro-Schiff bases and the asymmetric diimines from potentiometric data were carried out with the PKAS computer program (Martell and Motekaitis, 1988). The experimental method of Bjerrum and Calvin as modified by Irving and Rossotti (1954) was used to determine ň, ň_A and pL values. The stability constants of the complexes were calculated using the equation

2.6 Antimicrobial screening

In vitro antimicrobial activities of the asymmetric diimines (H₂L¹, H₂L² and H₂L³) and their Ni(II) and Fe(III) complexes were tested using the well-diffusion method (Nartop et al., 2008) against the bacteria: M. luteus (NRLL B-4375), B. cereus (RSKK 863), E. coli (ATCC 11230) in Molten Nutrient Agar medium and the fungus: C. albicans (ATCC 10239) in YPD Agar medium.

The bacterial strains were inoculated into Molten Nutrient Broth and incubated at 37 °C for 18 h. The yeast cells were inoculated into YPD Broth and incubated at 37 °C for 18 h. The test solutions of 10 mg mL⁻¹ and 4 mg mL⁻¹ in DMSO were prepared for all the diimines and the complexes, respectively. The YPD Agar (15 mL) kept at ca. 45 °C, was then poured in the petri-dishes which were inoculated with strains of yeast by taking 100 μL from cell culture media, and allowed to solidify. Then holes of 6 mm diameter were punched carefully using a sterile cork borer and these holes were filled with 50 μL of the test solutions. Petri-dishes were incubated for 24 h for bacteria and 48 h for fungi at 37 °C. The mean value obtained for the three holes was used to calculate the zone of growth inhibition of each sample. DMSO alone was used as a control under the same condition for each microorganism. MIC values (μg mL⁻¹) of H₂L¹ were done using serial dilutions (1, 2, 4, 8, 16, 32, 63, 125, 250, 500, 1000 μg/mL).

3.1 Ligand protonation constants

Stoichiometric protonation constant values and potentiometric titration curves of the nitro-Schiff bases (SB¹ and SB²) and the asymmetric diimines (H₂L¹, H₂L² and H₂L³) are presented in Figs. 1 and 2 respectively.

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Figure 1 The structures and the protonation constants of all the Schiff bases in 1:1 (v:v) ethanol–water solution at 25.0 ± 0.1 °C (^∗values in parentheses are standard deviations, and they are less than 0.01 for all of the constants).

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Figure 2 Potentiometric titration curves for the nitro-Schiff bases and the diimines.

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The nitro-Schiff bases have two protonation constants: first constant (log K_OH¹) refers to the protonation of phenolate anion for SB¹ and naphtholate anion for SB², second constant (logK_NH¹) belongs to the protonation of imine nitrogen for both. As can be seen in Fig. 2, the breaks near m = 1 for the SB¹ and SB² arise from releasing one hydrogen ion from the protonated form of the ligands. Both of the logK_OH¹ and logK_NH¹ values of SB² are found lower than that of the SB¹. Hence, in the case of SB², the electron densities of imine nitrogen and naphtholate oxygen atoms are smaller. It may be resulting from the higher electron withdrawing effect of the naphthalene ring compared with the benzene ring. This approach is in line with a literature report for similar compounds (Gündüz et al., 1993; Köseoğlu et al., 1995).

The asymmetric diimines consist of two different fragments bound to central 1,4-phenylene ring:

• On the left side, salicyl/naphthaldimine fragment attaches to the center by nitrogen atom. It may form a six membered chelate ring with intramolecular hydrogen bond.

• Phenyl/naphthyl azomethine moiety links by carbon atom to the center on the right side. It may also form five membered chelate ring with intramolecular hydrogen bond.

The asymmetric diimines have four protonation constants: first and second constants (log K_OH¹ and log K_OH²) correspond to the protonation of phenolate and naphtholate anions, third and fourth constants (log K_NH¹ and log K_NH²) refer to protonation of imine nitrogen atoms.

Protonated form of the diimines has four acidic hydrogen atoms and can be shown as [H₄L]²⁺. Two hydrogens are attached to the imine nitrogen atoms and they are neutralized in the pH region 3–5. The breaks near m = 2 for their titration curves are due to the appearing of two hydrogen ion from the protonated form of the ligands (in Fig. 2). Other hydrogen ions are bounded to oxygen atoms of the hydroxyl groups and are neutralized between pH 8 and 10. After first two protonations, molecules return to neutral H₂L forms and tautomerism occurs.

The phenol–imine and keto–amine tautomers are formed for the aromatic Schiff bases that have the hydroxyl group in ortho position to the imino group. If the hydrogen atom is completely transferred to the imino nitrogen atom, keto–amine tautomer is formed (Ünver et al., 2001). In the case of weak hydrogen bond between them, phenol–imine tautomer is protected (Ünver et al., 2006; Hökelek et al., 2000).

In our previous work (Güngör and Gürkan, 2010), we investigated the phenol–imine and keto–amine tautomerism of the diimines (H₂L¹, H₂L² and H₂L³) with spectroscopic methods. We concluded that, in methanol solution, only phenol–imine tautomer was formed on both sides of H₂L¹. On the other hand, for H₂L² and H₂L³, ketoamine tautomer was dominant (about 60%) on the left side, while phenol–imine tautomer was formed on the right side, as shown in Fig. 3.

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Figure 3 Tautomeric forms of the H₂L² and H₂L³.

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The effect of the tautomeric forms on the protonation constants can be interpreted as follows:

1. for H₂L¹, the phenol–imine tautomer only occurs on both left and right side of the molecule. The logK_OH¹ and logK_OH² values (9.08 and 8.17) are related with the protonation of the phenolic oxygen atoms.

2. for H₂L² and H₂L³, the phenol–imine tautomer forms on the right side of the molecules (see Fig. 3). The logK_OH¹ values (8.27 and 8.06) of these molecules refer to the protonation of the phenolic oxygen atoms. On the other hand, the keto–amine tautomer forms predominantly on the left side of the molecules. Hence, logK_OH² values (7.35 and 7.25) of these molecules are preferably related with the protonation of less acidic keto oxygen atoms as shown in Fig. 3.

3. for all of the diimines, logK_NH² value of the nitrogen atom on the left side is lower than that of on the right side (logK_NH¹) of the molecule. It indicates that the basicity of the nitrogen atom on the left side is less than the other side. It can be explained by the increasing probability of (O—H…N) intramolecular hydrogen bonding in six membered chelate ring.

3.2 Complex stability constants

At first, the complex forming reaction was investigated. For this purpose, the titrations of Ni(II) and Fe(III) ions to H₂L¹ ligand in 1: 1 molar ratios were performed. The data are shown in Fig. 4. The break near m = 4 for the systems shows the release of four hydrogen ions from the protonated form of the ligand when complexes are formed. Hence, complex forming reactions can be given as follows: $${[{\text{H}}_4\text{L}]}^{2+}+{\text{M}}^{\text{n}+}\leftrightarrow {\text{ML}}^{(\text{n}-2)}+4{\text{H}}^{+}$$

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Figure 4 Potentiometric titration curves of nickel(II) and iron(III) complexes of H₂L¹.

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The stability constants of the complexes are listed in Table 1. The constants increase in the order NiL² < NiL¹ < NiL³ for Ni(II) complexes while FeL¹ < FeL² < FeL³ for Fe(III) complexes. The stability constant of H₂L³ ligand with the two metal ions is the biggest. This may be attributed to increasing electron delocalization of two naphthalene rings in the H₂L³ molecule.

3.3 Structure of the complexes

The proposed molecular formulas of the complexes are given in Fig. 5. After the reaction of ONNO type tetradentate diimines with Ni(II) chloride, dimeric and binuclear [NiL]₂ type square planar complexes are formed for H₂L² and H₂L³, while [NiL(H₂O)₂]₂ type octahedral complex is formed for H₂L¹. On the other hand, reaction of diimines with Fe(III) chloride gives the octahedral complexes of the type [FeL(H₂O)Cl]₂·nH₂O (n = 0, 1).

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Figure 5 The suggested structures of the complexes.

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All of the complexes are powder; they are stable at room temperature, insoluble in water and slightly soluble in polar organic solvents. Their melting points are above 360 °C. The molar conductance values of the complexes are measured to be 0.63–41.96 Ω⁻¹ cm² mol⁻¹ in 10⁻³ M DMF or DMSO solutions. This fact indicates their nonelectrolytic nature (Geary, 1971).

3.3.1

3.3.1 Mass spectra

The mass spectra (ESI-positive) of the NiL¹ and FeL³ (see Fig. 6) show the characteristic molecular ion [M]⁺ peaks at the desired position for the dimeric structure, while NiL³, FeL² and FeL¹ give the [M+1]⁺, [M+4]⁺ and [M+6]⁺ peaks, respectively. The molecular ion peak of NiL² could not be determined in the spectrum. The biggest fragment ion at m/z = 583.2 for NiL² complex (see Fig. 7), loss of 261 mass unit from the dimeric [Ni(C₂₄H₁₆N₂O₂)]₂ unit, may suggest the expulsion of [C₁₈H₁₂NO + 3H]⁺ unit from the dimeric molecule. In the spectra of NiL¹, NiL² and NiL³, the highest intensity peaks occurred m/z = 214 (100%), m/z = 264 (100%) and m/z = 314 (100%). The peaks at m/z = 214, m/z = 264 and m/z = 314 may correspond to [C₁₃H₁₄N₂O]⁺, [C₁₈H₄N₂O]⁺ and [C₂₁H₁₈N₂O]⁺, respectively. In the spectra of FeL¹, FeL² and FeL³, the highest intensity peaks occurred m/z = 267 (100%), m/z = 395 (100%) and m/z = 445 (100%). The peaks at m/z = 267, m/z = 395 and m/z = 445 may correspond to [C₁₈H₇N₂O]⁺, [C₂₄H₇N₂OFe]⁺ and [C₂₈H₉N₂OFe]⁺, respectively.

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Figure 6 The fragments observed in mass spectra of FeL³.

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Figure 7 The fragments observed in mass spectra of NiL².

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3.3.4

3.3.4 Thermal studies

The TGA curves of the complexes are given in (Fig. 8(a–e)). Full thermoanalytical data of the complexes are given in Table 2. In TGA curves, NiL¹ decomposes in two steps. Mass loss in the first thermal reaction is observed at temperatures between 83 and 260 °C, corresponding to the four moles of coordinated water molecule in the complex. NiL² and NiL³ decomposed in one step. The endothermic peaks lying at ca 430–536 °C for all of the Ni(II) complexes correspond to melting. The residues at the end of the decomposition for all of the complexes are found to be NiO and the percent of metal ion is calculated.

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Figure 8 The TGA curves for (a) NiL¹; (b) FeL¹; (c) NiL²; (d) FeL²; (e) NiL³; (f) FeL³.

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Table 2 Thermal data of the complexes.

Complex	1. Thermal reaction decomposition of H2O molecules		2. Thermal reaction decomposition of the complex
	Temperature range ( °C)	% Mass loss found (Calcd.)	Temperature range ( °C)	% Residue (NiO/Fe2O3) found (Calcd.)
NiL1	83–260	8.80 (8.82)	432–800	17.30 (18.31)
FeL1	45–205	6.20 (6.25)	320–711	16.80 (18.48)
NiL2	–	–	430–800	16.40 (17.70)
FeL2	51–266	3.80 (3.81)	Indefinite initiation and termination temperatures	17.20 (16.88)
NiL3	–	–	419–726	16.10 (15.83)
FeL3	52–242	3.40 (3.44)	Indefinite initiation and termination temperatures	12.90 (15.27)

In TGA curves, all of the Fe(III) complexes decompose in two steps. The initial mass loss occurs at ca 45–266 °C range for complexes; this mass loss seems to be consistent with evaluation of the one mole hydration water molecule and/or two coordination water molecules in the complex. The endothermic peaks lying at ca 364–410 °C for FeL¹ and 380–443 °C for FeL² were due to melting. The residues at the end of the decomposition for all of the complexes are found to be Fe₂O₃ and the percent of metal ion is calculated.

3.3.5

3.3.5 Powder X-ray diffraction studies

The powder XRD patterns of all the complexes were recorded at 2θ = 0–70° range. The X-ray diffractograms, the ‘d’ values and 2θ angles of NiL² and FeL² complexes are given in Fig. 9 and listed in Table 3. The spacing between the planes in the atomic lattice (d) was calculated using Bragg’s formula as follows: $$n·\lambda =2·d·\sin \theta$$ where ‘n’ is an integer (n = 1), ‘λ’ is the wavelength (λ = 1.54059 Å) and ‘θ’ is the peak angle.

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Figure 9 Powder XRD pattern of complexes (a) NiCl₂·6H₂O; (b) NiL²; (c) FeCl₃·6H₂O; (d) FeL².

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Table 3 Powder X-ray diffraction data of complexes NiL² and FeL².

Complex	Peak	2θ	d (Å)	Complex	Peak	2θ	d (Å)
NiCl2·6H2O	1	5.0	17.66	NiL2	1	4.5	19.62
	2	5.8	15.23		2	5.6	15.77
	3	7.2	12.27		3	7.2	12.27
	4	8.7	10.16		4	9.1	9.71
	5	16.3	5.43		5	11.7	7.56
	6	18.4	4.82		6	16.4	5.40
	7	20.5	4.33		7	18.7	4.74
	8	25.0	3.56		8	24.8	3.59
	9	29.8	3.00		9	25.8	3.45
	10	30.4	2.94		10	28.0	3.18
	11	31.2	2.86		11	30.4	2.94
	12	32.3	2.77		12	32.1	2.79
	13	33.2	2.70		13	37.0	2.43
	14	34.1	2.63		14	37.6	2.39
	15	35.3	2.54		15	42.1	2.15
	16	37.2	2.42		16	44.8	2.02
	17	39.6	2.27		17	46.0	1.97
	18	42.0	2.15		18	46.4	1.96
	19	43.0	2.10
	20	44.0	2.06
	21	45.0	2.01
	22	46.5	1.95
	23	47.5	1.91
	24	49.1	1.85
	25	49.8	1.83
	26	51.6	1.77
	27	53.5	1.71
	1	4.3	20.50		1	3.1	28.48
	2	5.1	17.31		2	5.0	17.66
	3	7.8	11.33		3	5.2	16.98
	4	8.8	10.04		4	9.3	9.50
	5	9.1	9.71		5	11.1	7.97
	6	10.0	8.84		6	12.8	6.91
	7	13.9	6.37		7	13.0	6.80
FeCl3·6H2O	8	14.8	5.98	FeL2	8	16.2	5.47
	9	15.7	5.64		9	18.7	4.74
	10	17.6	5.04		10	19.2	4.62
	11	19.4	4.57		11	21.1	4.21
	12	22.0	4.04		12	22.1	4.02
	13	24.2	3.68		13	23.2	3.83
	14	25.8	3.45		14	26.0	3.42
	15	26.7	3.34		15	27.4	3.25
	16	28.3	3.15		16	28.3	3.15
	17	30.0	2.98		17	29.2	3.06
	18	32.4	2.76		18	30.3	2.95
	19	35.0	2.56		19	33.1	2.70
	20	37.2	2.42		20	35.3	2.54
	21	39.2	2.30		21	37.2	2.42
	22	41.7	2.16		22	39.2	2.30
	23	46.5	1.95		23	41.2	2.19
	24	51.8	1.76		24	42.0	2.15
	25	53.8	1.70		25	44.3	2.04
	26	59.1	1.56		26	54.2	1.69
					27	56.7	1.62
					28	59.6	1.55

The diffractogram of NiL² complex has recorded 18 reflections of 2θ ranging from 4.5° to 46.4°. The maximum recorded is at 9.1°. The observed sin θ value and d values obtained have been compared with NiCl₂·6H₂O. The diffractogram of FeL² complex showed 28 reflections between 2θ ranging from 3.1 to 59.6 with maxima peaks at 2θ = 9.3° and 27.4°. The observed sinθ value and d values obtained have been compared with FeCl₃·6H₂O.

All of the Ni(II) and Fe(III) complexes show well defined sharp peaks. This shows that the above complexes are crystalline in nature (Nair and Joseyphus, 2008). The line broadening of the crystalline diffraction peak in the Ni(II) complexes shows higher crystallinity than that of Fe(III) complexes by comparing the diffractograms.

3.4 Antimicrobial activities

The antimicrobial activity of the asymmetric diimines and their Ni(II) and Fe(III) complexes were tested against bacteria: M. luteus (NRLL B-4375), B. cereus (RSKK 863), E. coli (ATCC 11230) and the fungus: C. albicans (ATCC 10239). The diffusion agar technique was used to determine the antimicrobial activity (Çete et al., 2006). The results of antimicrobial screening of the compounds are summarized in Table 4.

As can be seen in Table 4, H₂L¹ diimine and its Ni(II) and Fe(III) complexes have an antimicrobial activity on all the studied microorganisms. Because of the biggest activity of H₂L¹ against all the studied microorganisms, its minimum inhibitory concentration (MIC; μg mL⁻¹) was determined. It shows high levels of antimicrobial activity against M. luteus (Figs. 10(a) and 11) and C. albicans. H₂L² diimine and its Fe(III) complex show antimicrobial effect on the microorganisms, whereas NiL² complex is inactive against B. cereus. While H₂L³ diimine exhibits activity against the microorganisms, NiL³ complex is inactive against B. cereus and E. coli; FeL³ complex does not show any effect against B. cereus. Table 4 shows that E. coli is the most influenced bacteria against all of the diimines and the diameter inhibition zones (34 mm) of H₂L² diimine are the biggest against M. luteus. Moreover, all of the diimines have an activity on C. albicans. In all bacteria, M. luteus is the most influenced bacteria against Ni complexes, in the 36–22 mm diameter inhibition zones ranges. Between the Ni(II) complexes, NiL¹ exhibits the most activity against the tested bacteria. For the Fe(III) complexes, FeL¹ has the biggest diameter zones (30 mm) against M. luteus (Fig. 9(a)). The fungicidal screening shows that C. albicans is the most influenced against NiL¹ complex, having the diameter zones of 16 mm and FeL³ complex shows quiet activity against this fungi.

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Figure 10 The diameter inhibition zones for (a) FeL¹ and (b) NiL³ against M. Luteus; (c) H₂L¹ and (d) H₂L³ against C. albicans; (e) H₂L² against E. coli.

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Figure 11 The minimum inhibitory concentration (MIC) value of H₂L¹ against M. Luteus (left to right: 1000, 500, 250, 125, 63, 32, 16, 8, 4, 2, 1 μg/mL).

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