Synthesis, spectroscopic characterisation, biological and DNA cleavage properties of complexes of nicotinamide

2.1 Materials and methods

All the chemicals used for the preparation of the ligands were of BDH quality, AR grade. Molar conductance of the complexes was measured using a Systronic conductivity bridge at room temperature in DMSO. Conductivity measurements (Ω⁻¹ cm² mol⁻¹) were carried out in DMF using a Tacussel conductivity bridge model. Magnetic susceptibilities were determined using a Guoy balance at room temperature (25 °C) with Hg[Co(SCN)₄] as standard. Perkin-Elmer PE 938 spectrophotometers were used to record the IR spectra using KBr pellets. UV–visible spectra were recorded in acetone and chloroform on a Perkin-Elmer model 550-S spectrometer. EPR spectra were recorded using Broker Ac 200 spectrophotometer. The electrochemical experiments were carried out using a Trace-Iab50 from radiometer which includes a polarographic analyser (Pol 150), a polarographic stand (MDE 150) and trace Master 5 software performed using a conventional three electrodes system. Potentials are expressed versus the Ag/AgCl (KCl mol l⁻¹) electrode separated from the test solution by a salt bridge containing the solvent/supporting electrolyte. A pre polished glassy carbon (GC) disc of 3 mm diameter (radiometer) was used as the working electrode and a platinum wire was the auxiliary electrode.

Solutions of CT DNA (calf-thymus DNA) in 50 mM NaCl/50 mM tris–HCl (pH = 7.2) gave a ratio of UV absorbance at 260 and 280 nm, A260/A280 of ∼1.8–19, indicating that DNA was sufficiently free of protein contamination. DNA concentration was determined by UV absorbance at 260 nm after 1:100 dilutions. The molar absorption coefficient was taken as 6600 M⁻¹ cm⁻¹. Stock solutions were kept at 4 °C and used within 4 days. Double distilled water was used to prepare the buffer.

2.2 Preparation of the complexes

Distilled ethanol (5.0 ml) and metal sulphates (0.50 m mol) are added to a 50 ml erlenmeyer flask fitted with a micro watch glass cover containing a magnetic stirring bar. Then the metal salt was dissolved in 1.0 mol of sodium sulphate and 2.0 m mol of nicotinamide. A large excess of ethanol is used, as it helps in the reaction completion. Heating and stirring the mixture was done for ∼1 h. As the nicotinamide combines with the metal forms complexes, while cooling, the mixture crystals begin to form as a crust at the surface of the reaction mixture (Swamy and Pola, 2008; Swamy et al., 2003; Goeta et al., 2000).

2.3 General properties

Cr(III), Mn(II), Fe(III), Co(II), Ni(II) and Cu(II) complexes are coloured but the cadmium complexes are colourless (Table 1). All the complexes are soluble in dimethyl formamide (DMF). Some of them are soluble only in nitrobenzene and acetone.

3.1 Conductance

The observed molar conductance values of the complex in DMF are in the range of 10–20 Ω⁻¹ cm² mol⁻¹ at room temperature. The value for 1:1 electrolyte in DMF is of the order of 65–90 Ω⁻¹ cm⁻¹ mol⁻¹. Hence from the conductivity measurement, it is concluded that the sulphate ions are covalently bonded to metal ions, which indicates that they act as ligands and not as simple ions. Based on the metal–ligand ratio calculated by the analytical data and the nature of the electrolytes given by the conductance measurements, compositions were assigned for the prepared complexes. From the magnetic and conductometric analysis it is predicted that the complexes may have the following structures [Cr₂(L₂)₃(L₁)₉], [Mn(L₂)(L₁)₅], [Fe₂(L₂)₃(L₁)₉], [Co₂(L₂)₃(L₁)₉], [Ni (L₂)(L₁)₅], [Cu(L₂)(L₁)₅] and [Cd(L₂)(L₁)₅].

3.2 Magnetic moments

The Cr(III) complex showed a magnetic moment of 3.17 BM which is slightly lower than the spin only value of 3.87 BM expected for the three unpaired electrons, which offer possibility of an octahedral geometry. In this high spin complexes, the magnetic moments of the Fe(III) are close to the spin only value of 5.61 BM, because the ground state (derived, from the ⁵S state of the free iron) has no orbital angular momentum and there is no effective mechanism for introducing any by coupling with excited states. In the low spin complexes, with $t_{2g}^5\text{,}{e}_g^0$ configurations may have considerable orbital contributions to their moments at room temperature (Chandra et al., 2007). The experimental magnetic moments of the prepared Fe(III) complex indicates the high spin (S = 5) octahedral d⁵-system. The magnetic moment values for Cu(II), Co(II) and Ni(II) complexes are shown in Table 2. The magnetic moment of Co(II) complexes are in the range of 4.90 BM indicating that the Co(II) complexes are typically high spin complexes and having octahedral structure. The Ni(II) complexes exhibit the magnetic moment values in the range 3.8 BM, indicating octahedral co-ordination of the ligands around Ni(II) ion. The Cu(II) complexes exhibit magnetic moment in the range of 2.11 BM suggesting distorted octahedral nature for these complexes. It is obvious that the metal complexes possess antiferromagnetic properties by slight intramolecular antiferromagnetic spin exchange interaction for binuclear chromium and iron, complexes with nicotinamide ligands. The absorption bands observed for the electronic spectra of the metal complexes also support the octahedral geometry (Turhan-Zitouni et al., 2001a,b).

3.3 Infrared spectral analysis

The IR spectra of ligand (L) with its octahedral complexes have been studied in order to characterise their structures. The IR spectra of the free ligand and its metal complexes were carried out in the 4000–400 cm⁻¹ range. The IR spectra of all metal complexes were interpreted by comparing the spectra with those of the free ligand, and the results are listed in Table 3. The comparison of the band positions of various vibrations are ascertained with good evidence. In the infrared spectrum of the ligand, the band at 1385–1400 cm⁻¹ is due to asymmetric C⚌N stretching vibration. The presence of δ(CONH₂) asymmetric stretching vibration is confirmed by the band at 1740 cm⁻¹ and the symmetrical stretching vibration is observed at 3350 cm⁻¹. In the infrared spectrum of metal sulphate complexes, the ν CN stretching vibration were observed at 1560–1610 cm⁻¹ and was due to coordination of the nitrogen from CN to the metal, stretching vibration for L reduced at the complex. The free nicotinamide ligand showed a strong peak at 1460 cm^–1 for L, which is characteristic of the imine ν (CN) group. The ν (C⚌C) stretching vibrations are affected upon complexation and are situated at a frequency significantly different than the free ligands. Coordination of the nicotinamide ligands to the metal centre through the nitrogen atom is expected to reduce the electron density in the methine and imine link and hence lower the ν (C⚌C) and ν (CN) absorption frequencies. The peak due to ν (C⚌C) is slightly shifted to lower frequencies and appears between 1576 and 1579 cm⁻¹, indicating the coordination of the imine nitrogen to the metal.

Similarly the other metal complexes also show that the stretching frequency for the CN and C⚌C are shifted to lower indicating that the metal is coordinated through the imine nitrogen atom. The peak at 1680 cm⁻¹corresponds to the asymmetric C⚌O stretching vibration and the symmetric O⚌CN stretching vibration observed at 1750 cm⁻¹ confirms that all the complexes, the ligand do not coordinate with CONH₂ nitrogen.

The band at 1680 cm⁻¹ which is assignable to nicotinamide I band arising mainly from the CN stretching vibration in free ligand is found to have no change in frequencies (1680–1685 cm⁻¹) in the metal complexes. The nicotinamide III band at 1532 cm⁻¹ due to a coupled C⚌O stretching mode move to higher wave numbers (1537–1532 cm⁻¹) compared to that of free ligands indicating the coordination of nicotinamide through the pyridine nitrogen. This results in the decreased CN bond order with concomitant increase in the C–O bond order. The bands due to υ (C–N–C) and υ (C–O–C) remain almost unchanged in all the complexes indicating that the nitrogen and oxygen of carbonyl moiety are not involved in binding. The free ligand showed a medium intensity band at 3210 cm⁻¹ assigned to ν NH vibrations, which has been observed in the 3203–3207 cm⁻¹ region for the complexes. It can be observed that there is no considerable shift in the ν NH vibrations in the case of the complexes compared to the ligands indicating non-involvement of amide NH function in the coordination. A strong intensity band observed at 1675 cm⁻¹ and a medium intensity band at 1630 cm⁻¹ indicate that the ligands are assigned to ν (C⚌O) and ν (CN) functions respectively. In the case of complexes the band due to ν C⚌O was observed in the 1670–1676 cm⁻¹ region, indicating its non-involvement in the complexation. The low frequency skeletal vibration due to M–N stretching provides direct evidence of the complexation. In the present investigation, bands are observed in the 479–417 cm⁻¹ region for ν M–N vibrations, respectively (Chandra and Gupta, 2001; Rai et al., 2005). The occurrence of bands at ∼1100 and ∼600 cm⁻¹ suggests the presence of free sulphate ion (Td) due to ν₃ and ν₄ mode of vibrations, respectively. However in complexes, an additional series of six bands appeared at ∼1110, ∼1040, ∼970, ∼645, ∼600 and ∼437 cm⁻¹ indicating the coordination of sulphate group in monodentate manner through oxygen atom, the symmetry being lowered to C₃V upon coordination.

3.4 Electronic spectral analysis

Electronic spectra of ligand L and their metal complexes in DMF solutions have been recorded in the 200–1100 nm range. The UV–Vis spectra of the ligand and metal complexes in DMF showed two to seven numbers of absorption bands between 268 and 734 nm (Table 4). The bands below 455 nm are mostly associated with intra ligand π → π^∗ and n → π^∗ transitions. In the electronic spectra of the ligand and their metal complexes, the presence of a wide range of bands is due to both π → π^∗, n → π^∗ and d–d transitions and also due to charge transfer transition arising from π electron interactions between the metal and ligand that involves either a metal-to-ligand or ligand-to-metal electron transfer. The absorption bands observed within the range of 331–397 nm in DMF are mostly due to the transition of n → π^∗ of imine group corresponding to the ligand or metal complexes.

3.5 EPR spectral analysis

3.5.1

3.5.1 EPR spectrum of Cu(II) complex

The EPR spectra of the Cu(II) complexes were recorded (Fig. 1) in a polycrystalline sample as well as in solution at room temperature in different frequencies. The g-values are calculated by using the expression, $$G=2.0023(1-4\lambda /10\text{Dq})\text{,}$$ where λ is the spin–orbit coupling constant for the metal ion.

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Figure 1

EPR spectrum of Cu(II) complex.

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The analysis of spectra gives the values for g_|| 2.22–2.27 and $g^{\perp }=2.09-2.10$ . The observed ‘g’ values for the complexes are less than 2.3 in agreement with the covalent character of the metal ligand bond. The trend $g_{||}>g^{\perp }>2.0023$ observed for the complexes indicates that the unpaired electron is localised in dx²–y² orbital of the Cu(II) ion and spectral features are characteristic of axial symmetry. Tetragonal elongated structure is confirmed for the aforesaid complex.

3.5.2

3.5.2 Chromium complex

The spin–orbit coupling constant from the free ion value 90 cm⁻¹ of Cr(III) complex is reduced from its metal ion and this can be employed as a measure of metal–ligand covalency (Fig. 2). The values indicate that the complexes under study have a substantial covalent character. The g-values were calculated and found in the range of 1.97–1.99, corresponding to 6-coordinated geometry. It is possible to define a covalency parameter analogous to the nephelauxetic parameter, which is the ratio of the spin–orbit coupling constant for the complex and the free Cr(III) ions.

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Figure 2

EPR spectrum of Cr(III) complex.

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3.6 Thermal decomposition

The TG curves of the metal–nicotinamide complexes suggest that modes of decomposition resemble each other closely. All these complexes follow the same pattern of thermal decomposition. The complexes remain almost unaffected up to ∼70 °C (Table 5). After this, a slight depression is observed up to ∼110 °C. The weight loss in this temperature range is equivalent to two water molecules in the complexes indicating them to be lattice water in conformity with our earlier observations from analytical and IR spectral investigations. The anhydrous complexes remain stable up to ∼300 °C and thereafter the complexes show rapid degradation presumably due to decomposition of organic constitution of the complexes as indicated by the steep fall in the percentage weight loss. The decomposition continues up to ∼800 °C and reaches to a stable product in each complex as indicated by the constituency in weight in the plateau of the thermogram. The decomposition temperature varies for different complexes as shown in Table 5.

There is no mass loss observed up to 185 °C for the Cr complex. The sharp endothermic DTA peaks at 123–183 °C show that the complex melts before decomposition. The thermo gravimetric curves of all the complexes show single stage decomposition wherein simultaneous decomposition and oxidation of the organic moiety and the volatilisation of chromium take place after 165–185 °C. There are broad endothermic peaks at 274 °C indicating slow decomposition of the complex. Due to the oxidation of the organic moiety two broad exothermic peaks appear between 290 °C and 355 °C for the complex. No such exothermic peaks were observed in the case of manganese complex. Due to volatilisation the weight of the final residue does not conform to any particular and product.

For Fe complex, there is no mass loss up to 181 °C, two endothermic peaks appeared at 80 and 120 °C, which are due to the decomposition and loss of coordinated water molecules. There afterwards, the complex shows single stage decomposition without giving any stable intermediate compound. A broad exothermic peak at 220 °C followed by an endothermic peak at 280 °C shows the oxidative process. The weight of the final residue almost corresponds to the formation of FeO.

3.7 Electro chemical studies

Cyclic voltammogram of cobalt complex recorded in 1.0 mM electrolyte containing DMSO is shown in Fig. 3a. The samples were scanned in the range from −1.9 to +1.9 V at the scan rate of 20 mV S⁻¹. During the cathodic scanning process, Cobalt complexes show a reduction peaks at 0.7 V due to the deposition of Co(II) on electrode surface and an oxidation peak at 1.05 V due to the stripping of cobalt from the surface. Also a small reduction peak was observed about at −0.85 V, which is mainly attributed to the adsorption of the double hydroxide at the cathode. Under these conditions, the pH value near the cathode seems to rise due to the proton consumption. Therefore, the hydroxide of metal ion may be produced in the vicinity of the electrode.

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Figure 3

Cyclic voltammogram of (a) 1 mM Co(II) complex (b) Cd(II) complex in 0.1 M TBATFB in DMSO on Pt electrode, V = 0.1 Vs⁻¹, (vs. Ag|Ag⁺ electrode).

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The cyclic voltammogram of Cd(II) complex recorded from 0.1 M TBATFB in DMSO showed in Fig. 3b. Potential varied from −2.0 to 1.0 V at the scan rate of 50 mV s⁻¹. During the cathodic scan, no reducible species present from 1.0 to −0.85 V, cathodic peak observed at −1.2 V due to the dissolution of cadmium from the electrolyte solution. During the reverse process, three oxidation peaks were observed (−0.7 and −0.2 V and a peak at 0.45 V) which is corresponding to the oxidation of cadmium in three phases.

Cyclic voltammogram of chromium complex recorded in electrolyte containing water is shown in Fig. 4a. The samples were scanned in the range from −1.0 to +1.0 V at the scan rate of 20 mV S⁻¹. During the cathodic scanning process, chromium nicotinamide complex shows a reduction peaks at −0.25 V and an oxidation peaks at 0.8 V due to the stripping of chromium from the surface. Under these conditions, the pH value near the cathode seems to rise due to the proton consumption. Therefore, the hydroxide of metal ion may be produced in the vicinity of the electrode. Cyclic voltammogram of Ferric complex is recorded in 1 mM electrolyte containing water is shown in Fig. 4b. The samples were scanned in the range from −1.0 to +1.0 V at the scan rate of 20 mV S⁻¹. During the scanning process, a reduction peak was observed at 0.25 V due to the dissolution of hydroxide ion and the deposition of ferrous ion on the base metal. From the figure, we can conclude that the good electrochemical performance of the electrode can be attributed to its higher electrical conductivity and facile ionic transportation that are offered by the nicotinamide additive. This metal complex also undergoes transformation between +2 and +3, and the transformation between these two oxidation states is highly reversible.

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Figure 4

Cyclic voltammogram of (a) 1 mM Cr(III) complex (b) Fe(III) complex in 1 mM Fe(III) in DMSO on Pt electrode, V = 0.1 Vs⁻¹, (vs. Ag|Ag⁺ electrode).

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3.8 Powder X-ray analysis

The XRD (Powder Pattern) of the complexes [Co₂(L₂)₃(L₁)₉] and [Ni₂(L₂)(L₁)₅] were indexed in X-ray diffractometer and the unit cell parameters have been calculated with the help of a computer from 2θ values (Fig. 5). The direct constant parameters like A, B, C, α, β, γ and V (volume) are given in Table 6. The density of the complexes has been determined by the floatation technique in a saturated solution of NaCl, KBr and benzene separately.

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Figure 5

The XRD (powder pattern) of the complexes (a) [Co₂(L₂)₃(L₁)₉] and (b) [Ni₂(L₂)(L₁)₅].

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The number of formula units per unit cell (n) was calculated from the relation $$n=\mathit{dNV}/M\text{,}$$ where d = density of the compound, N = Avogadro’s number, V = volume of the unit cell and M = molecular weight of the complex. The value of ‘n’ is found to be 10 for both complexes which agree well with the suggested monoclinic structure of the complexes. In addition, we have carried out powder X-ray diffraction studies of complex. Powder XRD pattern of [Co(L₂)₃(L₁)₉] consists of 7 reflections in the range 5–50° (2θ) the inter planar spacing (d) has been calculated from the positions of intense peaks using Bragg’s relationship. The 2θ values with maximum intensity of the peak for ligand were found to be 5.709 (2θ) which correspond to d = 15.4539. All the important peaks have been indexed and the observed values of inter planar distance were compared with the calculated ones. It was found that there is good agreement between the calculated and observed values. The (h² + k² + l²) values are 1, 5, 27, 32, 58, 72 and 74. The presence of forbidden number confirms the tetragonal systems. This implies that the cobalt and nickel complexes are distorted octahedral in structure.

3.9 Antimicrobial activity

The biological and medicinal potency of the coordination complexes has been established by their antitumor, antiviral and antimalarial activities. This characteristic property has been related to the ability of the metal ion to form complexes with ligand containing nitrogen donor atoms. The ligands and their complexes were screened for their antibacterial activity (Plakatouras et al., 1992) against Escherichia coli and Staphylococcus aurious and antifungal activity against (Assour, 1965) Aspergillus niger and Aspergillus flavus at the concentration of 100 μg/0.1 cm³. The standard drugs streptomycin and clotrimazole were also tested for their antibacterial and antifungal activity at the same concentration under the conditions similar to that of the test compounds concentration.

3.10 Antibacterial activity

For in vitro antimicrobial activity, the investigated compounds were tested against the bacteria such as Shigella dysenteriae, E. coli and Bacillus subtilis. The minimum inhibitory concentration (MIC) values of the compounds against the growth of microorganisms are summarised in Table 7. It is observed that the copper and cadmium complexes are more active in Pseudomonas aeruginosa and B. subtilis, respectively compared to other bacterial organisms. Nickel and cobalt complexes are moderately active in all bacterial organisms compared with standard streptomycin. From the results it is found that the copper complex is more active in Streptococcus-β-haemolyticus than the other complexes.

3.11 Antifungal activity studies

The results of the antifungal screening of the nicotinamide and the metal complexes with Candida albicans, A. niger and Aspergillus fumigates at concentration of 200 μg by disc method are given in the Fig. 6. Comparative studies of the ligands and their complexes indicated that metal complexes exhibit higher antifungal activity than the free ligands (Table 8). The antifungal activity results revealed that the ligands and their Cu(II), Co(II) and Ni(II), complexes have exhibited weak to good activity against A. niger and A. flavus. The ligand and its Cu(II) and Co(II) complexes show weak activity when compared to the standard drug clotrimazole. The order of the metal complexes follow Cu(II) > Cd(II) > Ni(II) > Co(III) > Mn(II) > Fe(III) > Cr(III).

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Figure 6

Anti fungal activity of Aspergillus flavus of (A) ligand (B) [Ni(L₂)(L₁)₅].

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The higher activity of metal complexes can be explained on the basis of overtons concept and chelation theory. According to overtons concept of cell permeability, the lipid membranes that surround the cell favour the passage of only the lipid soluble material due to which lip solubility is an important factor, which controls antimicrobial activity. On chelation, the polarity of metal ion will be reduced to a greater extent due to the overlap of the ligand orbital and partial sharing of the positive charge of the metal ion with donor group. Further it increases the delocalisation of pi-electron over the whole chelate ring, lyophilizes, enhances the penetration of the complexes into lipid membranes blacking the metal binding sites in the enzymes of microorganisms. These complexes also disturb the respiration process of the cell and thus block the synthesis of protein that restricts further growth of the organism.

3.12 DNA cleavage studies

The metal complexes were able to convert super coiled DNA into open circular DNA. The general oxidative mechanisms proposed account for DNA cleavage by hydroxyl radicals. The general oxidative mechanisms proposed account for the abstraction of a hydrogen atom from sugar units predicting the release of specific residues arising from transformed sugars, depending on the position from which the hydrogen atom is removed. The cleavage is inhibited by the free radical scavengers implying that hydroxyl radical or peroxy derivatives mediate the cleavage reaction. The reaction is modulated by a metallo complex bound hydroxyl radical or a peroxo species generated from the co-reactant H₂O₂.

In the present study, the CT–DNA gel electrophoresis experiment was conducted at 35 °C using our synthesized complexes in the presence of H₂O₂ as an oxidant. It was found that, at very low concentrations, few complexes exhibit nuclease activity in the presence of H₂O₂. Control experiment using DNA alone does not show any significant cleavage of CT-DNA even on longer exposure time. Hence, we conclude that the copper complex cleaves DNA as compared with control DNA, while other complexes do not cleave DNA in the presence of H₂O₂. Probably this may be due to the formation of redox couple of the metal ions and its behaviour.

The redox property of the metal complexes mediates oxidation of nucleic acids. In oxidative DNA cleavage mechanism, metal ions in the complexes react with H₂O₂ to generate the hydroxyl radical which attacks at the C4′ position of the sugar moiety and finally cleaves the DNA. Copper complex reacts with H₂O₂ to produce hydroxyl radical, hydroxyl ion and Cu(II) form. The formation of hydroxyl radical by the copper complex is further compared with other complexes with H₂O₂. Hence, copper complex can promote redox mediated cleavage of DNA reaction on sugar ring. The Cu(II) ion formation is supported by electrochemical study. Further, the presence of a smear in the gel diagram indicates the presence of radical cleavage.

As can be seen from the results (Fig. 7), at a very low concentration, few complexes exhibit nuclease activity in the presence of H₂O₂. Control experiment using DNA alone (lane 1) does not show any significant cleavage of CT-DNA even on longer exposure time. From the observed results, we conclude that the complexes, copper complex (lane 6), nickel complex (lane 7) and cobalt complex (lane 8) cleave DNA as compared to control DNA while other complexes (lanes 1–5) do not cleave DNA in the presence of H₂O₂. Probably this may be due to the formation of redox couple of the metal ions and its behaviour. Further, the presence of a smear in the gel diagram indicates the presence of radical cleavage.

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Figure 7

DNA cleavage studies of ligand and metal complexes.

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