DFT calculations, spectroscopic studies, thermal analysis and biological activity of supramolecular Schiff base complexes

2.1 Analysis and physical measurements

Carbon, hydrogen and nitrogen were analyzed by standard microanalysis methods using Automatic analyzer CHNS Vario EL III-Elementar, Germany at Microanalytical center, Cairo University, Giza, Egypt. Mass spectra of the complexes were performed using GC-2010 Shimadzu at Microanalytical center, Cairo University, Giza, Egypt. FTIR spectra were obtained with KBr disk technique using test scan Shimadzu FTIR spectrometer. The spectra were collected in the range 400–4000 cm⁻¹. Molar conductivity of 10⁻³ mol/L solutions of the complexes in DMF was measured on the conductivity meter ORION model 150 of 0.6 cell constant. Thermal analysis of Mn₂—L and Co—HL complexes was carried out using TG-50 Shimadzu Thermal Analyzer from ambient temperature to 1000 °C at a heating rate 10 °C min⁻¹ in dry nitrogen atmosphere at a flow rate of 30 mL min⁻¹ in platinum crucible. The molar Magnetic susceptibility of the chelates was measured at room temperature on prepared samples using a magnetic susceptibility Cambridge England Sherwood Scientific. The effective magnetic moments were calculated from the expression μ_eff = 2.828(X_M · T)^1/2 B.M., where X_M is the molar susceptibility corrected using Parcel’s constants for diamagnetism of all atoms in the compounds (Emeleus and Sharpe, 1974). The electronic absorption spectra in DMF were measured using automated UV/Vis–NIR 3101 PC Shimadzu spectrophotometer ranged from 200 to 900 nm. The metal content in the prepared chelates was determined by the following wet digestion technique. This is performed by weighing 20–25 mg of the metal chelate, and adding 5 mL of concentrated nitric acid, followed by 5 mL hydrogen peroxide. The solution was heated gently to fumes and then boiled until the solution becomes clear, excess of H₂O₂ was added if necessary then the solution was diluted with little double distilled water and neutralized with sodium hydroxide solution to pH 5.5 and the metal content was determined by the recommended method (Macdonald and Sirichanya, 1969).

2.2 Synthesis of metal complexes

The solid complexes were prepared by the addition of a hot ethanolic solution containing one mmol of Cu(CH₃COO)₂·H₂O (199 mg), Ni(ClO₄)₂·6H₂O (366 mg), Co(CH₃COO)₂·6H₂O (285 mg) or MnCl₂·4H₂O (198 mg) to 10 mL ethanolic solution containing one mmol (540 mg) of the Schiff base H₂L, m.p. 300 °C (Literature m.p. is 300 °C, (Amin et al., 2012)), Scheme 1, which was synthesized as previously described (Amin et al., 2012). The solution was stirred on a water bath for 2 h during which the metal complex precipitated. The resulting complexes were filtered out and washed thoroughly with successive portions of hot ethanol followed by petroleum ether until the filtrate becomes colorless. The obtained complexes were kept in a vacuum desiccator over anhydrous calcium chloride. The complexes are air stable in the solid state and soluble in DMF or DMSO.

2.3 Biological activity measurement

The antibacterial activities of the Schiff base and its complexes were investigated using the conventional disk-agar diffusion method (El-Gaby, 2004). The activity of the tested samples was studied against the Staphylococcus aureus and Bacillus subtilis as gram-positive bacteria (G⁺) and Pseudomonas aeruginosa and Escherichia coli as gram-negative bacteria (G⁻) at 20 mg/mL concentration. The bacterial strains were cultured on nutrient agar media. The broth media of bacteria were incubated for 24 h with subsequent filtering of the culture through a thin layer of sterile glass G2. 50 mL of agar media was seeded with 1.00 mL of bacteria and poured into Petri dishes and cooled down to room temperature.

The tested compounds were dissolved in DMSO (which has no inhibition activity) with different concentrations. Sterile uniform size 8 mm diameter filter paper disks (3 disks per compound) were dipped into the dissolved tested compounds and carefully placed on surface of the seeded agar plate. After incubation for 24 h at 37 °C, the diameter of inhibition zone surrounding each disk was measured and used to calculate the mean of inhibition zones. Ampicillin was used as antibacterial standard agent.

2.4 Molecular orbital calculations

In this section, we try to explore the optimized geometrical parameters (bond lengths, bond angles and dihedral angles), net charges on active centers and energetic of the ground state for the ligand (H₂L) and its complexes with Mn(II), Co(II), Ni(II) and Cu(II).

In all calculations, Density Functional Theory (DFT) used at the B3LYP level (Becke, 1993), 3-21G (McGrath and Radom, 1991) as a basis set for the ligand and LANL2DZ (Hay and Wadt, 1985) as basis set for the complexes. All calculations were performed using G09 W program (Frisch et al., 2009).

3.1 Mass spectra of metal complexes

The mass spectra were measured in order to confirm the composition and to support the identity of the structure of the prepared metal complexes, Fig. 1. The mass spectra of mononuclear complexes revealed the molecular ion peak at m/z 772 (calcd. 722.30) and 686 (calcd. 686.51) for [Ni·HL·OH⁻·3H₂O]·3H₂O and [Co·HL·OH⁻·3H₂O]·H₂O, respectively. In case of binuclear complexes their molecular ion peeks are in agreement with their calculated formulae [Mn₂·L·2Cl⁻·6H₂O]·H₂O and [Cu₂·L·2OH⁻·2H₂O]·2H₂O at m/z 845 (calcd. 845.39) and 771 (calcd. 771.67), respectively.

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Figure 1

Mass spectra of the metal complexes.

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3.2 FTIR spectra

Table 2 shows the infrared band assignments of Schiff base (H₂L) and its metal chelates.

IR spectra of the complexes were compared with that of the free Schiff base in order to show the bonding sites in the ligand that may be involved in the chelation. The discussion was focused on the most important bands ν (C⚌O)_benzopyrone, ν (C⚌N)_azomethine, ν (OH) and ν (C—O)_phenolic in the IR spectrum of the free Schiff base.

The metal complexes contained all the bands from the ligand and also other bands indicative of the coordination of the ligand with the metal ions.

The carbonyl stretching vibration ν (C⚌O)_benzopyrone of two benzopyran-4-one moieties appeared as two bands in the range 1675–1658 cm⁻¹ in the IR spectrum of the ligand (Amin et al., 2012). These bands shifted to lower wave number after complexation due to the perturbing effect of the coordinated metal ions. The shift in binuclear complexes is different from that of mononuclear complexes. In binuclear complexes, the band appears as doublet at lower wave number, Table 2, while in mononuclear complexes, the ν (C⚌O)_benzopyrone appears as two bands. This means that the coordination of metal ion to —C⚌N— in one side affects the carbonyl of this side, while ν (C⚌O)_benzopyrone of the unchelated side is slightly changed. However, the band due to ν (C⚌N)_azomethine at 1636 cm⁻¹ in the spectrum of the free ligand shifted to lower frequency (43–54 cm⁻¹) and interfered with the band of stretching frequency of C⚌C in all the metal complexes. This indicates the participation of this group in the complexation (Bradley et al., 1981). This is supported by the appearance of new band in the spectra of all complexes at 423–488 cm⁻¹ which is attributed to νM—N stretching (Bradley et al., 1981). The shift, Table 2, occurs in the band at 1167 cm⁻¹ (Amin et al., 2012) due to ν (C—O)_phenolic refers to the coordination through the oxygen of the hydroxyl group in position seven of the benzopyran-4-one moiety and the formation of C—O—M bond (Bellamy, 1975; Saraswat et al., 1977). In the low frequency region, new band observed in the FTIR spectra of all chelates in the region 511–542 cm⁻¹ may be attributed to νM—O (Rao, 1963). This in turn, confirms the bond formation between the metal ion and the oxygen of phenolic group of the benzopyran-4-one ring.

The broad band at 3420–3422 cm⁻¹ in the IR spectra of all complexes could be attributed to the stretching vibrations of water molecules (Junnosuke et al., 1956; Kemp, 1978). The band at 876–886 cm⁻¹ in the FTIR spectra of the complexes is due to the OH bending of the coordinated water molecules. This confirms the presence of the coordinated water molecules (Nakamoto, 1986).

3.3 Electronic absorption spectra and Magnetic susceptibility measurements

The electronic absorption spectra of the metal complexes in dimethylformamide (DMF) are shown in Fig. 2. The electronic absorption spectrum of Schiff base H₂L as discussed before (Amin et al., 2012) exhibits a band at 390 nm (25,641 cm⁻¹) which is due to an intramolecular charge transfer (CT transition) involving the whole molecule. Another shoulder appears at 340 nm (29,412 cm⁻¹) in the spectrum of free Schiff base which can be assigned due to the π–π^∗ transition within the azomethine (CH⚌N) chromophore group (Downing and Urbach, 1969). A bathochromic shift of charge transfer band is observed in the electronic spectra of Cu₂—L and Mn₂—L complexes at 27,778 and 28,985 cm⁻¹, respectively, while hypsochromic shift is observed in the spectra of Ni—HL and Co—HL complexes at 30,675 and 30,581 cm⁻¹, respectively. This shift confirms the participation of nitrogen atom of azomethine group in the complexation (Sönmez and Sekerci, 2002). The electronic spectra of all complexes show bands in the range of 23,041–26,316 cm⁻¹, Table 3, which may be mainly due to LMCT transition. The electronic spectrum of Ni—HL complex shows an additional band at 25,773 cm⁻¹ attributed to ³A_2g → ³T_1g(P) transition (Patil et al., 2013) which indicates an octahedral geometry around Ni(II) ion.

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Figure 2

Electronic absorption spectra of Schiff base and its complexes.

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The magnetic moment value of the binuclear complex (2:1) (M:L) Cu₂—L complex is 2.76 B.M (1.38 B.M. per metal ion) and this μ_eff is consistent with the proposed square planar geometry for the Cu(II) complex (Chohan and Jaffery, 2000). On the other hand, the magnetic moment of Ni—HL complex was (3.33) B.M. This value is within the range (2.9–3.3 BM), found for paramagnetic complex of Ni(II) with octahedral geometry (Nicholls, 1973; Jha and Mukharjee, 2007). The room temperature magnetic moment of the solid cobalt (II) complex lie at 4.37 B.M., indicating the presence of three unpaired electrons (Earnshaw, 1968) in high spin octahedral environment (Chavan et al., 2011; Gudasi et al., 2006) which is in good agreement with the octahedral range 4.3–5.2 BM (Hankare et al., 2004a,b). In addition, the binuclear Mn₂—L complex shows high spin paramagnetic moment of 10.14 B. M. (5.07 B.M. per metal ion) corresponding to the presence of five unpaired electrons in an octahedral environment (Thakor et al., 2010).

3.4 Thermogravimetric analysis

Thermogravimetric analyses data, Table 4, were measured from ambient temperature to 1000 °C at a heating rate of 10 °C/min. The TG/DTG curves of Co—HL and Mn₂—L complexes are shown in Fig. 3. The first decomposition step was observed in the TG curve of Co—HL complex within the temperature range of 23–100 °C appeared as peak at 50 °C in DTG curve corresponds to the loss of one hydration water molecule with mass loss of 2.73% (calcd. 2.62%). The second step of decomposition in the region 100–178 °C corresponds to the loss of three coordinated water molecules with mass loss of 7.68% (calcd. 7.87%). This step appeared as peak at 157 °C in the DTG curve of this complex. The loss of organic moiety occurs at the temperature range of 178–819 centered at 475 °C. The complex completes its gradual decomposition till the stable cobalt oxide is formed as a final product above 819 °C.

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Figure 3

TG/DTG curves of Mn₂—L and Co—HL complexes.

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The thermal decomposition for this complex can be formulated as follows: $$[\text{Co}·\text{HL}·{\text{OH}}^{-}·3{\text{H}}_2\text{O}]·{\text{H}}_2\text{O}\stackrel{23''–''100{}^{°}\text{C}}{\to }[\text{Co}·\text{HL}·{\text{OH}}^{-}·3{\text{H}}_2\text{O}]$$ $$[\text{Co}·\text{HL}·{\text{OH}}^{-}·3{\text{H}}_2\text{O}]\stackrel{100''–''178{}^{°}\text{C}}{\to }[\text{Co}·\text{HL}·{\text{OH}}^{-}]$$ $$[\text{Co}·\text{HL}·{\text{OH}}^{-}]\underset{178''–''819{}^{°}\text{C}}{\overset{\text{Loss of organic part}}{\to }}\text{CoO}$$

On the other hand, the loss of one lattice water molecule, mass loss of 2.08% (calcd. 2.13%), occurs at the temperature range of 20–72 °C and appeared as peak at 51 °C in the TG/DTG curve of Mn₂—L complex. The mass loss of 12.92% (calcd. 12.78%) in the range 72–159 °C, DTG peak at 100 °C, is due to the elimination of six coordinated water molecules from this complex. The chloride ion was eliminated in the 159–243 °C range (Gaber and Al-Shihry, 2004) with mass loss 8.33, calcd. (8.39%). The organic part decomposes in the range 243–820 °C with DTG peak at 514 °C to give the manganese carbonate, MnCO₃ as a final product.

On this basis, the thermal decomposition for Mn₂—L complex can be formulated as follows: $$[{\text{Mn}}_2·\text{L}·2{\text{Cl}}^{-}·6{\text{H}}_2\text{O}]{\text{H}}_2\text{O}\stackrel{20''–''{72}^{°}\text{C}}{\to }[{\text{Mn}}_2·\text{L}·{\text{Cl}}^{-}·6{\text{H}}_2\text{O}]$$ $$[{\text{Mn}}_2·\text{L}·2{\text{Cl}}^{-}·6{\text{H}}_2\text{O}]\stackrel{72''–''{159}^{°}\text{C}}{\to }[{\text{Mn}}_2·\text{L}·2{\text{Cl}}^{-}]$$ $$[{\text{Mn}}_2·\text{L}·2{\text{Cl}}^{-}]\stackrel{159''–''{243}^{°}\text{C}}{\to }[{\text{Mn}}_2·\text{L}]$$ $$[{\text{Mn}}_2·\text{L}]\underset{243''–''820{}^{°}\text{C}}{\overset{\text{Loss of organic part}}{\to }}[{\text{MnCO}}_3]$$

According to elemental analyses data, interpretation of IR spectra, thermal analyses, conductivity, electronic spectra and magnetic susceptibility measurements, it may be concluded that the metal ions coordinated to the Schiff base (HL⁻) via the azomethine nitrogen atom and the oxygen of the deprotonated hydroxyl group in case of (1:1) (M:L) Ni—HL and Co—HL complexes, so the ligand behaves as a monobasic bidentate, Scheme 2. On the other hand, Cu(II) and Mn(II) ions were bonded to L²⁻ ligand through the two C⚌N groups and the two oxygen of the deprotonated hydroxyl groups to form binuclear (2:1) (M:L) complexes in which the ligand acts as dibasic tetradentate, Scheme 3.

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Scheme 2

The proposed structures of (a) Ni—HL and (b) Co—HL complexes.

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Scheme 3

The proposed structures of (a) Cu₂—L and (b) Mn₂—L complexes.

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3.5 Geometrical parameters of the ligand

The optimized structure of the ligand is shown in Fig. 4 with numbering of the atoms. The energies of the ground state and the dipole moment computed by B3LYP/3-21G for H₂L are presented in Table 5.

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Figure 4

The optimized geometry, vector of dipole moment and numbering system for H₂L using B3LYP/3-21G.

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The computed energy gap E_g (E_LUMO − E_HOMO) is an important stability indicator that helps to characterize the chemical reactivity and kinetic stability of the molecule (Rakha et al., 2014). As the E_g decreases the reactivity increases and the amount of electronic charge transfer from the ligand to the central metal ion increases i.e. it easily offers electrons to an acceptor. Low value of energy gap is also due to the groups that enter into conjugation (Rakha et al., 2014). Accordingly, the ligand L²⁻ is more reactive than HL⁻, as reflected from the energy gap values, by 0.215 eV (5.8 kcal). Thus, L²⁻ ligand can act as a tetradentate ligand and interacts with two metal ions to form binuclear metal complexes in case of Cu₂—L and Mn₂—L. From the computed net charge on the active centers, Table 6, noting that the most negative centers are O7, O35, N9 and N17 hence, the ligand bonded through these active centers.

3.6 Geometrical parameters of the complexes

Tables 5–8 give the values of the energies, dipole moment, energy gap, energy of HOMO and LUMO, net charge on active centers, bond lengths and bond angles of all the complexes under study.

In addition, Figs. 5 and 6 represent the optimized geometry, numbering system and the vector of the dipole moment for metal complexes using B3LYP/LANL2DZ. In complexes of Co—HL and Ni—HL, the metal coordinates with N9 and O7 while in Cu₂—L and Mn₂—L with N9, O7, N17 and O35.

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Figure 5

The optimized geometry, numbering system and vector of the dipole moment for Ni—HL and Co—HL complexes using B3LYP/LANL2DZ.

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Figure 6

The optimized geometry, numbering system and vector of the dipole moment for Mn₂—L and Cu₂—L complexes using B3LYP/LANL2DZ.

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The geometric changes that are observed in the ligand moiety itself are interesting. Thus, most of bonds show elongation upon complexation with the metal ions. From the analysis of the data calculated for the bond lengths, Table 7, one can conclude that C8—N9, C3—O7, C24—N17 and C29—O35 bond lengths become longer in all complexes, as the coordination takes place via N atoms of the (C⚌N)_azomethine and the oxygen of the deprotonated hydroxyl groups (El-Gammal, 2010). This finding is due to the formation of the M—O and M—N bonds which make the C—O and C—N bonds weaker (El-Gammal, 2010). The bonds between M and the ligand sites in the complexes i.e. M—N and M—O are long, 1.874 Å for Ni—N9, 1.800 Å for Ni—O7, 1.924 Å for Co—N9, 1.846 Å for Co—O7, 1.862 Å for Cu—N9, 1.921 Å for Cu—O7, 1.902 Å for Mn-N9 2.151 Å for Mn—O7, 1.850 Å for Cu—N17, 1.931 Å for Cu—O35, 2.018 Å for Mn—N17 and 2.215 Å for Mn—O35 compared to the typical M—X bond lengths which have the values Cu—O 1.62 Å, Cu—N 1.66 Å and Ni—N 1.44 Å (Armelao et al., 2010). The long M—O and M—N bonds in the complexes mean that the ionic character of these bonds is small. Also, the charge on the metal ion in the complex is much less than 2 in Co(II), Ni(II), Cu(II) and Mn(II). Hence, the comparison between the calculated B3LYP/LANL2DZ and the typical M(II)—N and M(II)—O is not very precise. The bond angles of the Schiff base moiety are altered relatively upon coordination but the angles around the metal vary upon changing the metal center. This change affects the other angles, Table 8, which are reduced or increased on complex formation as a consequence of bonding. The bond angles confirm that the Ni—HL, Co—HL and Mn₂—L complexes are distorted octahedral with two water molecules occupying the axial positions and the other water molecule and the OH⁻ or Cl⁻, required to complete the coordination number of the metal, at the equatorial positions to give octahedral geometry, while Cu₂—L complex shows square planar geometry.

The values of the dihedral angles around metal ion in the complexes are far from 0.0° or 180° which indicate that the metal ion is not in the same plane of the donating sites. All the studied complexes are polar as it is evident from the magnitude of their dipole moments, Table 5. As the energy gap of the studied complexes decreases, the reactivity of the complexes increases. The reactivity of the complexes under study follows the order Cu₂—L > Mn₂—L > Ni—HL > Co—HL.

Table 6 presents the net charge on active centers of the ligand and the studied metal complexes and the amount of charge transferred from the ligand to the central metal ions i.e. L → M. The results of Table 6 show that in Ni—HL and Co—HL complexes, Ni12 and Co12 ends up with a net charge of 0.342 e and 0.294 e, respectively, indicating that the two metal ions received 1.658 e and 1.706 e from their surrounding donating sites of the ligand (O7 and N9). In case of Cu₂—L and Mn₂—L the charge density increases on the donating atoms after complexation indicating that there is an electron back–donation from the metal ion to the donating sites. These findings are further confirmed by comparing the values of the calculated charge on the donating sites in both the complexes and the free ligand.

So theoretical calculations confirm the results obtained from the above mentioned tools of analysis that the metal ions coordinated to the Schiff base (HL⁻) via the azomethine nitrogen atom (N9) and the oxygen of the deprotonated hydroxyl group (O7) in case of (1:1) (M:L) Ni—HL and Co—HL complexes and the ligand behaves as a monobasic bidentate, Fig. 5. For binuclear (2:1) (M:L) complexes one metal is coordinated with N9 and O7 and the other metal ion with N17 and O35 indicating that the ligand acts as dibasic tetradentate, Fig. 6.

3.7 Biological activity

The main aim of the production and synthesis of any antimicrobial compound is to inhibit the causal microbe without any side effects on the patients. In addition, it is worthy to stress on the basic idea of applying any chemotherapeutic agent which depends essentially on the specific control of only one biological function and not multiple ones.

In testing the antibacterial activity of the compounds under investigation we used more than one tested organism to increase the chance of detecting antibiotic principles in tested compounds. The antibacterial activity of the Schiff base (H₂L) and its metal complexes against S. aureus and B. subtilis (G⁺) bacteria and P. aeruginosa and E. coli as (G⁻) bacteria were tested, Fig. 7, to inhibit the causal microbe without any side effects on the patients.

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Figure 7

The antibacterial activities of Schiff base and its complexes.

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The biological activity of the ligand and its metal complexes were also compared with Ampicillin (Antibacterial agent) as a standard. In the present study the ligand, Ni—HL and Cu₂—L showed no biological activity towards these types of bacteria, and in other words, they gave no inhibition zone diameters values. The two complexes Co—HL and Mn₂—L showed a biological activity towards S. aureus and B. Subtilis (G⁺) and P. aeruginosa and E. coli (G⁻) bacteria with inhibition zone diameters ranging from 11–15 and 20–22 mm/mg sample, respectively, Fig. 7. The Mn₂—L Complex has the highest antimicrobial activity among the complexes, almost close to that of the standard, Table 9.

The lipid membrane that surrounds the cell favors the passage of only the lipid-soluble materials, so lipophilicity is an important factor controlling antibacterial activity. This finding suggests that chelation can facilitate the ability of such complexes to cross the cell membrane and can be explained by Tweedy’s Chelation Theory (Dharmaraj et al., 2001). On chelation, the polarity of the metal ion will be reduced because of partial sharing of its positive charge with donor groups and the delocalization of π-electrons over the whole chelate ring will be increased. This enhances the lipophilicity of the complex which favors its penetration through the lipid membrane and blocks the metal binding sites in the enzymes of microorganisms.

In other words, the chelates may disturb the respiration process of the microbial cells. This will lead to the inability of the microbial cells to synthesize their own proteins that restrict further growth of the organism. Accordingly, this demonstrates the enhancement of antibacterial activity as a function of chelation. The variation in the activity values of different compounds against different organisms depends on either the impermeability of the cells of the microbes or on the differences in ribosomal structures of microbial cells (Dros et al., 1998). On the other hand, the Ni—HL and Cu₂—L complexes have no biological activities. This may be due to lower permeability of these complexes through the microbial cell membrane indicating the decrease of their penetration through the lipid membrane. Hence, they could neither block the enzyme’s active site nor inhibit the growth of the microorganisms.

The biological activity of the prepared complexes can be correlated with ground state properties. In spite of, the higher reactivity of Cu₂—L complex which can be explained in terms of the energy gap, it has no biological activity. It is found that Mn₂—L complex has higher antimicrobial activity than that of the free ligand and the other complexes. Mn₂—L complex has the highest dipole moment among the complexes as shown in Table 5. This finding may be one of the reasons for the highest biological activity of this complex.