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Journal of Chemistry
Volume 2013 (2013), Article ID 653540, 11 pages
http://dx.doi.org/10.1155/2013/653540
Research Article

Synthesis, Spectral, Thermogravimetric, XRD, Molecular Modelling and Potential Antibacterial Studies of Dimeric Complexes with Bis Bidentate ON–NO Donor Azo Dye Ligands

1P.G. Department of Chemistry, G.M. Autonomous College, Sambalpur, Odisha 768 004, India
2Department of Chemistry, Government College of Engineering, Kalahandi, Bhawanipatna, Odisha 766 002, India

Received 14 May 2013; Accepted 23 September 2013

Academic Editor: Mallikarjuna Nadagouda

Copyright © 2013 Bipin Bihari Mahapatra et al. This is an open access article distributed under the Creative Commons Attribution License, which permits unrestricted use, distribution, and reproduction in any medium, provided the original work is properly cited.

Abstract

The dimeric complexes of Co(II), Ni(II), Cu(II), Zn(II), Cd(II), and Hg(II) with two new symmetrical ON–NO donor bis bidentate (tetradentate) azo dye ligands, LH2 = 4,4′-bis(4′-hydroxyquinolinolinylazo)diphenylsulphone, and L′H2 = 4,4′-bis(acetoacetanilideazo)diphenylsulphone have been synthesized. The metal complexes have been characterised by elemental analytical, conductance, magnetic susceptibility, IR, electronic spectra, ESR, NMR, thermogravimetry, X-ray diffraction (powder pattern) spectra, and molecular modelling studies. The Co(II) and Ni(II) complexes are found to be octahedral, Cu(II) complexes are distorted octahedral, and a tetrahedral stereochemistry has been assigned to Zn(II), Cd(II), and Hg(II) complexes. The thermogravimetric study indicates that compounds are quite stable. The energy optimized structures are proposed using the semiempirical ZINDO/1 quantum mechanical calculations. The potential antibacterial study of the ligands and some metal complexes has been made with one gram positive bacteria Staphylococcus aureus and one gram negative bacteria E. coli which gives encouraging results. Both the Co(II) complexes are found to possess monoclinic crystal system.

1. Introduction

Azo dyes are found to be pharmacologically active and hence they are used as chemotherapeutic agents in the manufacture of potential drugs [1, 2]. Azo dyes are also used as indicator in the chemical laboratories and as preservative and dyeing agents in food industries [3]. Besides their applications azo dyes can also form stable complexes with transitional and nontransitional metal ions because of the presence of azo (–N=N–) group [49]. The present study reports the synthesis of two new ON–NO donor tetradentate azo dyes and their twelve dimeric complexes and characterization of these complexes are made by using various physicochemical methods. Antibacterial study and thermogravimetric and molecular modelling of the azo dyes and some metal complexes have also been described.

2. Experimental

2.1. Materials

The chemicals 4,4′-diaminodiphenyl sulphone, 8-hydroxyquinoline, and acetoacetanilide were E. Merck grade. The chlorides of Co(II), Ni(II), Cu(II), Zn(II), Cd(II), and Hg(II) were of S.R.L. grade. All other reagents and solvents were purchased from commercial sources and were of analytical grade.

2.2. Synthesis of the Azo Dye Ligands

The azo dyes were synthesized by the coupling reaction of the diazonium chloride obtained from 4,4′-diaminodiphenyl sulphone (0.01 mol, 2.0 g) with alkaline solution of 8-hydroxyquinoline (0.02 mol, 2.9 g) and acetoacetanilide (0.02 mol, 3.54 g) separately at 0–5°C.

2.3. Synthesis of Metal Complexes

The metal chlorides in ethanol were mixed separately with ethanolic solution of the ligands in 2 : 1 molar ratio and the resulting solutions were heated to 60–70°C for about one hour on a heating mantle. The solution was then cooled down to room temperature and the pH was raised to ~7 by adding conc. ammonia drop by drop with stirring. The solid complexes thus separated were then washed with ethanol followed by ether and dried in vacuum.

2.4.  Physical Measurements

The elemental analysis (C, H, N) were carried out on elemental analyser Perkin-Elmer 2400, while metals were determined by EDTA after decomposing the complexes with conc. HNO3. Conductance measurements of the complexes were made using Toshniwal CL 01–06 Conductivity Bridge. The magnetic susceptibility were made at RT by Gouy method using Hg[Co(CN)4] as calibrant. IR spectra (KBr) were recorded using IFS 66U spectrophotometer, electronic spectra of the CoII, NiII, and CuII complexes in DMF were recorded on a Hilger-Watt uvispeck spectrophotometer, ESR spectra of CuII complex were recorded on a E4-spectrometer, and NMR spectra were recorded on a Jeol GSX 400 with DMSO as solvent and TMS as internal standard. X-ray diffraction (powder pattern) of the CoII complexes was recorded on a Phillips PW 1130/00 diffractometer and the TG, DTG, and DTA of the complexes were recorded on NETZSCH STA 409 C/CD in nitrogen atmosphere at a heating rate of 10°C per minute.

2.5. Molecular Modelling

Molecular modelling of the ligand and complexes has been made by Argus Lab 4.0.1.

2.6. Antibacterial Activity

The antibacterial activity of the azodye ligands and CoII, NiII, CuII, and ZnII complexes were studied as per cup-plate method [10] using two strains of bacteria like Staphylococcus aureus and E. coli. The solutions of the test compounds are prepared in dimethylsulfoxide (DMSO) at 500 µg/mL. The bacterial strains are inoculated into 100 mL of the sterile nutrient broth and incubated at °C for 24 hours. The density of the bacterial suspension is standardized by McFarland method. A well of uniform diameter (6 mm) is made on agar plates after inoculating them separately with the test organisms aseptically. The standard drug and the test compounds are introduced with the help of micropipette and the plates are placed in the refrigerator at 8–10°C for proper diffusion of drug into the media. After two hours of cold incubation, the petri plates are transferred to incubator and maintained at °C for 18–24 hours. Then the petri plates are observed for zone of inhibition by using vernier scale. The results are reported by comparing the zone of inhibition shown by the test compounds with standard drug tetracycline. The results are the mean value of zone of inhibition of three sets measured in millimetre.

3. Results and Discussion

The physical characteristics and microanalytical data of the ligands and the complexes are given in (Table 1). The analytical data of the complexes revealed 2 : 1 molar ratio (metal : ligand) which corresponds well with the general formula [M2L/L′Cl2(H2O)6] and [L/L′Cl2(H2O)2] where M = CoII,  NiII, CuII; M′ = ZnII, CdII, HgII; LH2 = 4,4′-bis(4′-hydroxyquinolinylazo)diphenylsulphone, C30H20N6O4S (Calcd. (%) C, 64.28; H, 3.6; N, 14.99; Found (%) C, 63.8; H, 3.3; N, 14.5). L′H2 = 4,4′-bis (acetoacetalidoazo)diphenylsulphone, C32H30N6O6S (Calculated (%) C, 61.33; H, 4.83; N, 13.41; Found (%) C, 61.1; H, 4.6; N, 13.2). All the complexes are amorphous in nature and have high melting points and are insoluble in common organic solvents like methanol, ethanol, and benzene but soluble in dimethylformamide and dimethylsulfoxide. Nonelectrolytic nature of the complexes is indicated from the low conductance values (4.2–5.6 Ω−1 cm2 mol−1) in 10−3 M solution in DMF [11].

tab1
Table 1: Analytical and physical data of the ligands and its complexes.
3.1. IR Spectra

In the IR spectra of the azodye ligands (Table 2) a broad band obtained at 3390 cm−1 (LH2) and at 3443 cm−1 (L′H2) be assigned to O–HN and O–HO intramolecular hydrogen bonding. The absence of this band in the spectra of metal complexes indicates the deprotonation of hydrogen bonded NH or OH group on complexation and subsequent coordination of the phenolic/enolic oxygen atoms to the metal ions [12]. The sharp band of the ligands appear at 1625 cm−1 (LH2) and at 1633 cm−1 (L′H2) can be attributed to ν(–N=N–) vibration. There is no shift of this band in the metal complexes indicating noncoordination of the azo group to the metal ions. The band observed at 1148 cm−1 (LH2) is attributed to ν(C–O) vibration and the bathochromic shift of ~15 cm−1 in the metal complexes indicates bonding of oxine oxygen to the metal ions [13]. In the spectrum of the ligand (LH2) an intense band is observed at 1407 cm−1 due to C–N vibration of the oxinate group [14]. In the metal complexes this band occurs at ~1324 cm−1. The shift of this band to lower frequency regions shows considerably lower double bond character of the C–N bond due to involvement of the ring nitrogen on complexation [15, 16]. In the ligand (L′H2) the band observed at 1668 cm−1 can be assigned to ν(C=O) vibration and shifting of this band by 20–25 cm−1 to lower frequency region in the metal chelates indicates the coordination of the amidic oxygen atoms to the metal ions. The band shown at 1263 cm−1 in the ligand (L′H2) can be assigned to enolic (C–O) vibration and decrease of this frequency by 20–30 cm−1 on complexation is indicative of bonding of enolic oxygen atoms to the metal ions. In the metal complexes broad bands appear at ~3350–3399 cm−1 followed by sharp peaks at ~833–842 cm−1 and at ~727–736 cm−1 assignable to –OH starching, rocking, and wagging vibrations, respectively, indicating the presence of coordinated water molecules in the complexes [17]. The conclusive evidence of bonding of the azo dye ligands to the metal ions is proved by the appearance of bands at ~508–514 cm−1  ν(M–O) and ~450–460 cm−1  ν(M–N) [18].

tab2
Table 2: Infrared spectra of the ligand and the complexes in .
3.2. Electronic Spectra and Magnetic Measurements

In the electronic spectrum of CoII complexes, four ligand field bands are observed at 8200(8250), 16400(16500), 19730(19960), and 31545(32450) cm−1. The first three bands can be attributed to (ν1), (ν2), and (ν3) transitions, respectively, and the fourth band is assigned to a CT band. The ligand field parameters like  cm−1,  cm−1,  cm−1, ν2/ν1 = 2(2), and (24.37)% suggest an octahedral stereochemistry for the CoII complexes [19]. In the electronic spectra of NiII complexes, four ligand field bands are observed at 10115(10140), 16930(17125), 24825(24975), and 31345(32165) cm−1 assignable to (ν1), (ν2), (ν3), and CT transition, respectively, in an octahedral geometry. The ligand field parameters like  cm−1,  cm−1,  cm−1, ν2/ν1 = 1.673(1.688), and % also confirm an octahedral symmetry for the complexes [20]. The electronic spectra of the copper(II) complexes exhibit one broad band at 13300–14470 cm−1 with maxima at 13320(13345) cm−1 assignable to transition in support of a distorted-octahedral configuration of the copper (II) complex [21, 22]. The magnetic moment of the metal complexes were recorded at RT. The observed magnetic moment value of the CoII,  NiII, and CuII complexes are found to be ~5.0, ~3.1, and ~1.8 B.M., respectively, indicating octahedral configuration of the complexes, which is further supported by their electronic spectral data [23, 24].

3.3. 1H-NMR Studies

The 1H NMR spectra of the ligands LH2 and L′H2 were recorded in DMSO. The complex pattern observed at 6.746–9.344 ppm and at 7.039–7.956 ppm corresponds to eighteen phenyl protons in each ligand. The sharp peak obtained at 13.629 ppm LH2 corresponds to two phenolic protons. The sharp peaks obtained at 3.570 ppm, at 2.507 ppm, 10.913 ppm, and at 13.026 ppm in the ligand L′H2 correspond to six methyl (–CH3) protons, two methylene (>CH) protons, two amino (>NH) protons, and two enolic (>C–OH) protons, respectively [25]. (Figures 5(a) and 5(b)).

3.4. ESR Studies

The ESR spectra of the CuII Complexes [Cu2LCl2(H2O)6] and [Cu2L′Cl2(H2O)6] have been recorded at X-band at RT. The “gav” values of the complexes are found to be 2.09623 and 2.08807, respectively, by applying Kneubuhl’s method [26]. This type of spectrum may be due to dynamic or pseudorotational type of Jahn-Teller distortion (Figures 6(a) and 6(b)). The spin-orbit coupling constant () can be calculated from the equation The value of the former complex is found to be −320.445 cm−1 and that of latter complex is −293.823 cm−1. The decrease of the values of the complexes from the free ion value (−830 cm−1) indicates the overlapping of metal-ligand orbitals in the metal complexes.

3.5. Powder XRD Studies

The XRD study (powder pattern) of the complexes [Co2LCl2(H2O)6] and [Co2L′Cl2(H2O)6] has been made with the help of X-ray diffractometer with Cu as anode material, K-alpha [nm] = 0.154060, and the generator settings 30 mA, 40 KV. The prominent peaks of X-ray diffraction pattern have been indexed and analysed by using computer programme from LSUCRPC [27]. The lattices parameters like , , , , , , and (volume) are shown in (Tables 4(a) and 4(b)) along with miller indices . The indexing is confirmed by comparing between observed and calculated () values. It is observed that the peaks of the XRD powder pattern (Figures 7(a) and 7(b)) that have successfully indexed as figure of merit () is found to be 6.9 and 8.8, respectively, as suggested by de Woulff [28]. The density () of the complex was determined by the floatation method in a saturated solution of KBr, NaCl, and benzene separately. The number of formula units per unit cell () is calculated from the relation where = density of the compound, = Avogadro’s number, = volume of the unit cell, and = molecular weight of the complex. The value of “” is found to be 2 in both cases which agrees well with the suggested structure of the complexes. The crystal system of both the complexes was found to be monoclinic. The Debye-Scherrer equation in X-ray diffraction and crystallography is a formula which relates the size of the crystallites in a solid to the broadening of a peak in a diffraction pattern. The Debye-Scherrer equation is where = crystallite size, = wavelength of X-ray radiation (CuKα = 0.154060 nm), = constant taken as 0.94, = diffraction angle (23.08)°, and = full width at half maximum height (FWHM) (2.52 nm). The crystallite size of the complex [CO2LCl2(H2O)6] is found to be 4.99 nm. For the other complex [CO2L′Cl2(H2O)6], = diffraction angle (18.29)°, and = full width at half maximum height (FWHM) (2.77 nm). So crystallite size of this complex is found to be 2.61 nm [29].

3.6. Thermogravimetric Study

The complex [Ni2L′Cl2(H2O)6] suffers a mass loss of 3.4% at 100°C which corresponds to the removal of two lattice held H2O molecules supported by an endothermic peak on the DTA curve at 95°C [30]. Again, the complex moiety loses a mass of 23.52% at 250°C indicating removal of all coordinated H2O molecules and 1/6th of the ligand mass supported by an endothermic peak at about 240°C on the DTA curve. Thereafter at 450°C compound loses a mass of 23.07% which corresponds to the removal of 1/3rd of the ligand moiety supported by an exothermic peak at 420°C. Again the compound loses a mass of 37.5% indicating removal of 2/3rd of the ligand moiety. Again the compound loses 55% mass which corresponds to the removal of rest of the ligand moiety and two chlorine atoms and formation of NiO as residue (Figure 8(a)). The complex [Co2L′Cl2(H2O)6] loses a mass of 11.6% at 150°C with the removal of all coordinated H2O molecules supported by an endothermic peak at 140°C on the DTA curve. Then, the compound loses a mass of 13.5% indicating removal of 1/6th of the ligand moiety supported by an endothermic peak at 240°C. Thereafter, the complex moiety suffers a mass loss of 15.15% at 400°C which corresponds to the removal of 1/5th of the ligand moiety supported by an endothermic peak at 380°C. Finally the compound loses a mass of 64% at 700°C indicating removal of rest of the ligand moiety and two chlorine atoms with the formation of CoO as the residue (Figure 8(b)). The complex [Ni2LCl2(H2O)6] suffers a mass loss of 23.52% at 150°C indicating removal of all the coordinated H2O molecules along with 1/6th of the ligand supported by an endothermic peak at 140°C on the DTA curve. Then the compound loses a mass of 24.24% at 400°C which corresponds to the removal of 1/3rd of the ligand moiety supported by an exothermic peak at 325°C on the DTA curve. Finally, the compound loses 64% of mass indicating removal of rest of the ligand moiety and two chlorine atoms which is supported by an endothermic peak at 930°C on the DTA curve with the formation of NiO as the residue (Figure 8(c)).

The kinetic parameters such as order of reaction and activation energy for the thermal decomposition of [Cu2L′Cl2(H2O)6], [Ni2LCl2(H2O)6], and [Ni2L′Cl2(H2O)6] complexes have been determined by Freeman-caroll [31] method. In this method, the equation used is where = rate of heating, = weight fraction of reacting materials, = activation energy, = order of reaction, and = frequency. This equation in the difference form will be ; when is kept constant, a plot at versus Δlog  will give a linear relationship whose slope and intercept provide the value of and , respectively. The order of the decomposition reaction, the activation energy, and correlation coefficient are given in (Table 5). The calculated values of the activation energy is found to be low due to the autocatalytic [32, 33] effect of the metal ion on the thermal decomposition of the complex.

3.7. Optimized Geometry Studies of the Ligands & Complexes by Molecular Modelling Method

Molecular modelling of the ligands (LH2), (L′H2) and metal complexes of Co(II) have been carried out using molecular mechanics and Hartree-Fock (HF) Quantum methods. The standard 6–31 g basic set was used in conjugation with the HF method. All calculations are made using Gaussian 98 programme package [3437].

The metal complexes were built and the optimization of their geometries was done at mm/H–F/6–31 g level of theory Figures 1, 2, 3, and 4. The findings of these computed works are in good agreement with the experimental results. The selected bond lengths, bond angles of the ligand, bond angles of the complexes, and their bond energies are given in Tables 3(a), 3(b), 3(c), 3(d), 3(e), 3(f), 3(g), and 3(h), respectively. The total energies of both the complexes have been found to be 287.403 kcal/mole and 247.322 kcal/mole, respectively.

tab3
Table 3: (a) Selected bond lengths and bond energies of the ligand (LH2). (b) Selected bond angles and bond energies of the ligand (LH2). (c) Selected bond lengths and bond energies of the ligand (L′H2). (d) Selected bond angles and bond energies of the ligand (L′H2). (e) Selected bond lengths and bond energies of the [Co2LCl2(H2O)6] complex. (f) Selected bond angles and bond energies of the [Co2LCl2(H2O)6] complex. (g) Selected bond lengths and bond energies of the [Co2L′Cl2(H2O)6] complex. (h) Selected bond angles and bond energies of the [Co2L′Cl2(H2O)6] complex.
tab4
Table 4: (a) X-ray diffraction data of the complex [Co2LCl2(H2O)6]. (b) X-ray diffraction data of the complex [Co2L′Cl2(H2O)6].
tab5
Table 5: Kinetic parameters of the complexes.
653540.fig.001
Figure 1: Optimised geometry of ligand (LH2).
653540.fig.002
Figure 2: Optimised geometry of ligand (L′H2).
653540.fig.003
Figure 3: Optimised geometry of [Co2LCl2(H2O)6] complex.
653540.fig.004
Figure 4: Optimised geometry of [Co2L′Cl2(H2O)6] complex.
fig5
Figure 5: (a) 1H NMR spectra of LH2. (b) 1H NMR spectra of L′H2.
fig6
Figure 6: (a) ESR spectra of the [Cu2LCl2(H2O)6] complex. (b) ESR spectra of the [Cu2L′Cl2(H2O)6] complex.
fig7
Figure 7: (a) XRD graph for [Co2LCl2(H2O)6] complex. (b) XRD graph for [Co2L′Cl2(H2O)6] complex.
fig8
Figure 8: (a) TG/DTA graph of [Ni2L′Cl2(H2O)6] complex. (b) TG/DTA graph of [Co2L′Cl2(H2O)6] complex. (c) TG/DTA graph of [Ni2LCl2(H2O)6] complex.
3.8. Antibacterial Activity

The ligands and metal complexes have been screened for antibacterial activities and results have been shown in (Table 6). The antibacterial activity of the compounds is examined against two strains of bacteria, one gram positive Staphylococcus aureus and one gram negative E. coli. The effectiveness of the compounds is classified into three categories, Sensitive, intermediate, and resistant. If a compound is sensitive to a bacteria then it can be applied to cure the disease caused by the bacteria, while it fails to do so if it is resistant to the bacteria. Accordingly the effectiveness of the compound can be predicted by knowing the zone of inhibition value in mm. The results (Figure 9) show that the ligand was found to posses more antibacterial activity than the complexes against different bacteria. The increase in biological activity of the metal complexes than the ligands may be due to complexation and it can be explained on the basis of chelation theory [38].

tab6
Table 6: Antibacterial activities of the ligands and the complexes (data presented as diameter of zone of inhibition, mm).
653540.fig.009
Figure 9: Effect of the complexes on the growth of selected E. coli and S. aureus.

4. Conclusion

The CoII and NiII complexes are found to be octahedral and CuII complexes distorted-octahedral; ZnII, CdII, and HgII complexes are assigned to have tetrahedral geometry. Both the azo dyes behave as dibasic tetradentate ligands coordinating through oxine nitrogen, phenolic oxygen, enolic oxygen, and amidic oxygen atoms. All the complexes are dimeric in nature. The complexes are found to be thermally stable. From the thermal study of the complexes the order of decomposition reaction, activation energy and correlation coefficients has been calculated. The XRD study indicates a monoclinic crystal system for both the CoII complexes. All calculations based on molecular mechanics on the optimized geometries fit well with the experimental findings. The crystallite sizes of the complex compounds have been determined. The potential antibacterial study of the ligands as well as CoII, NiII, CuII, and ZnII complexes has been made against gram positive and gram negative bacteria which gives encouraging results.

Acknowledgments

The authors are thankful to The Head, SAIF, and I.I.T. Madras, India, for providing spectral analysis, MMIT, Bhubaneswar, for kind help of XRD data, and Dr. J. Panda, Department of Microbiology, Roland Institute of Pharmacy, Berhampur, Odisha, India, for providing antibacterial data.

References

  1. L. S. Goodman and A. Gilman, The Pharmacological Basis of Therapentics, Macmillan, New York, NY, USA, 4th edition, 1970.
  2. K. N. Gaind and J. M. Khanna, Indian Journal of Pharmaceutical Sciences, vol. 26, p. 34, 1949.
  3. R. M. Isa, A. K. Ghoneium, H. A. Dessouki, and M. M. Mustafa, “Co(II), Ni(II) and Cu(II) complexes of some phenylazosalisylaldehyde derivatives,” Journal of the Indian Chemical Society, vol. 61, pp. 286–289, 1984.
  4. B. B. Mahapatra, R. R. Mishra, and A. K. Sarangi, “Synthesis, characterisation, XRD, molecular modelling and potential antibacterial studies of Co(II), Ni(II), Cu(II), Zn(II), Cd(II) and Hg(II) complexes with bidentate azodye ligand,” Journal of Saudi Chemical Society, 2013.
  5. B. B. Mahapatra and S. K. Panda, “Coordination compounds of CoII, NiII, CuII, ZnII, CdII and HgII with tridentate ONS donor azo dye ligands,” Biokemistri, vol. 22, no. 2, pp. 71–75, 2011.
  6. B. B. Mahapatra and S. K. Panda, “Polymetallic complexes,” Indian Journal of Chemistry, vol. 87, pp. 1447–1452, 2010.
  7. B. B. Mahapatra and S. K. Panda, “Polymetallic complexes. Part-XCIX: tetrameric and dimeric CoII, NiII, CuII, ZnII, CdII and HgII complexes with hexa- and tetradentate azodye ligands,” Indian Journal of Chemistry, vol. 87, pp. 1199–1204, 2010.
  8. B. B. Mahapatra, A. K. Sarangi, S. K. Panda et al., “Polymetallic complexes part C dimeric Co(II), Ni(II), Cu(II), Zn(II), Cd(II) and Hg(II) complexes with bis-bidentate azodye ligands,” Jtr Chemicals Corporation, vol. 16, no. 2, pp. 59–63, 2009.
  9. B. B. Mahapatra and A. K. Sarangi, “Polymetallic complexes. Part-LXIV: hexadentate O O N–N O O donor azodye tetrameric complexes of CoII, NiII, CuII, ZnII, CdII and HgII,” Journal of the Indian Chemical Society, vol. 86, pp. 559–563, 2009.
  10. R. S. Brandt and E. R. Miller, “Studies with the agar cup-plate method: I. A standardized agar cup-plate technique,” Journal of Bacteriology, vol. 38, no. 5, pp. 525–537, 1939.
  11. J. V. Quagliano, J. Fujita, G. Franz, D. J. Phillips, J. A. Walmsley, and S. Y. Tyree, “The donor properties of pyridine N-oxide,” Journal of the American Chemical Society, vol. 83, no. 18, pp. 3770–3773, 1961. View at Scopus
  12. F. A. Cotton and P. G. Wilkinson, Advanced Inorganic Chemistry, Wiley Eastern, New Delhi, India, 3rd edition, 1985.
  13. L. K. Mishra and B. K. Keshari, “Thiohydrazides as complexing agent part 1-complexes of Ni(II), Co(II & III), Cu(II), Zn(II), Cd(II), Pd(II) & Hg(II) with O-Hydroxyphenylthiohydrazide,” Indian Journal of Chemistry A, vol. 28, pp. 883–887, 1981.
  14. P. B. Dorian, H. H. Patterson, and P. C. Jordan, “Optical spectra of Os4+ in single cubic crystals at 4.2°K,” Journal of Chemical Physics, vol. 49, no. 9, p. 3845, 1968. View at Publisher · View at Google Scholar
  15. R. Magee and L. Gordan, “The infrared spectra of chelate compounds-I: a study of some metal chelate compounds of 8-hydroxyquinoline in the region 625 to 5000 cm−1,” Talanta, vol. 10, no. 8, pp. 851–859, 1963. View at Publisher · View at Google Scholar
  16. R. K. Bajaj, G. S. Sodhi, and N. K. Kashia, “Halide and complex halogeno anions as salts of oxinato bis(η5-indenyl)titanium(IV)/zirconium(IV) chelates,” Polyhedron, vol. 3, no. 7, pp. 883–887, 1984. View at Publisher · View at Google Scholar
  17. G. S. Sodhi, A. K. Sharma, and N. K. Kaushik, “Halide and complex halogeno anions as salts of oxinate chelates of titanium(IV),” Journal of Organometallic Chemistry, vol. 238, no. 2, pp. 177–183, 1982. View at Publisher · View at Google Scholar
  18. K. Nakamoto, Infrared Spectra of Inorganic and Co-Ordination Compounds, Wiley Interscience, New York, NY, USA, 1983.
  19. J. R. Ferraro, Low Frequency Vibration of Inorganic and Coordination Compounds, Plenum Press, New York, NY, USA, 1971.
  20. A. B. P. Lever, Electronic Spectroscopy, Elsevier, Amsterdam, The Netherlands, 1968.
  21. A. B. P. Lever, “The electronic spectra of tetragonal metal complexes analysis and significance,” Coordination Chemistry Reviews, vol. 3, no. 2, pp. 119–140, 1968. View at Publisher · View at Google Scholar
  22. C. R. Hare and C. J. Ballahusen, “Crystal spectrum and magnetism of Tetrakis-Thiourea-Nickel Chloride,” Journal of Chemical Physics, vol. 40, p. 788, 1984. View at Publisher · View at Google Scholar
  23. S. Yamada, “Recent aspects of the stereochemistry of schiff-base-metal complexes,” Coordination Chemistry Reviews, vol. 1, no. 4, pp. 415–437, 1966. View at Publisher · View at Google Scholar
  24. C. K. Jorgensen, “Comparative crystal field studies. II. Nickel(II) and copper(II) complexes with polydentate ligands and the behaviour of the residual places for co-ordination,” Acta Chemica Scandinavica, vol. 10, pp. 887–910, 1966. View at Publisher · View at Google Scholar
  25. D. H. Williams and I. Fleming, Spectroscopic Methods in Organic Chemisty, Tata McGraw-Hill, Chennai, India, 4th Edn edition, 1994.
  26. F. K. Kneubuhl, “Line shapes of electron paramagnetic resonance signals produced by powders, glasses, and viscous liquids,” Journal of Chemical Physics, vol. 33, p. 1074, 1960. View at Publisher · View at Google Scholar
  27. J. M. Visser, “A fully automated programme for finding the unit cell from power data,” Journal of Applied Crystallography, vol. 2, no. 3, pp. 89–95, 1969. View at Publisher · View at Google Scholar · View at Scopus
  28. P. M. De Woulff, “A simplified criterion for the reliability of a powder pattern indexing,” Journal of Applied Crystallography, vol. 1, pp. 108–113, 1968. View at Publisher · View at Google Scholar
  29. A. Patterson, “The Scherrer formula for X-ray particle size determination,” Physical Review, vol. 56, no. 10, pp. 978–982, 1939. View at Publisher · View at Google Scholar
  30. A. Abu-Hussen, “Synthesis and spectroscopic studies on ternary bis-Schiff-base complexes having oxygen and/or nitrogen donors,” Journal of Coordination Chemistry, vol. 59, no. 2, pp. 157–176, 2006. View at Publisher · View at Google Scholar
  31. E. S. Freeman and B. Carrol, “The application of thermoanalytical techniques to reaction kinetics: the thermogravimetric evaluation of the kinetics of the decomposition of calcium oxalate monohydrate,” Journal of Physical Chemistry, vol. 62, no. 4, pp. 394–397, 1958. View at Publisher · View at Google Scholar
  32. A. M. El-Award, “Catalytic effect of some chromites on the thermal decomposition of KClO4. Mechanistic and non-isothermal kinetic studies,” Journal of Thermal Analysis and Calorimetry, vol. 61, p. 197, 2000.
  33. A. Impura, Y. Inoue, and I. Yasumori, “Catalysis by “Copper Chromite”. I. The effect of hydrogen reduction on the composition, structure, and catalytic activity for methanol decomposition,” Bulletin of the Chemical Society of Japan, vol. 56, no. 8, pp. 2203–2207, 1983. View at Publisher · View at Google Scholar
  34. M. A. Thomson and M. C. Zerner, “A theoretical examination of the electronic structure and spectroscopy of the photosynthetic reaction center from Rhodopseudomonas viridis,” Journal of the American Chemical Society, vol. 113, no. 22, pp. 8210–8215, 1991. View at Publisher · View at Google Scholar
  35. A. K. Rappé and W. A. Goddard III, “Charge equilibration for molecular dynamics simulations,” Journal of Physical Chemistry, vol. 95, no. 8, pp. 3358–3363, 1991. View at Scopus
  36. A. K. Rappe, K. S. Colwel, and J. Cassewit, “Application of a universal force field to metal complexes,” Journal of Inorganic Chemistry, vol. 32, no. 16, pp. 3438–3450, 1993. View at Publisher · View at Google Scholar
  37. J. Cassewit, K. S. Colwel, and A. K. Rappe, “Application of a universal force field to main group compounds,” Journal of the American Chemical Society, vol. 114, no. 25, pp. 10046–10053, 1992. View at Publisher · View at Google Scholar
  38. K. Mahanan and S. N. Devi, “Synthesis, characterization, thermal stability, reactivity, and antimicrobial properties of some novel lanthanide(III) complexes of 2-(N-salicylideneamino)-3-carboxyethyl-4,5,6,7- tetrahydrobenzo[b]thiophene,” Russian Journal of CoordinationChemistry, vol. 32, p. 600, 2006.