Synthesis, Characterization, and Antibacterial Activity of some Chelates in O,N-Donor Coordination Pattern involving Schiff bases derived from           4-Acetyl/Benzoyl-1-(4’-Nitrophenyl)-3-Methyl-2-Pyrazolin-5-one and Sulfamoxole

 

Thakar A.S.¹*, Friedrich H.B.¹ and Joshi K.T.²

¹School of Chemistry, University of KwaZulu-Natal, Durban, 4000, South Africa
²Department of Chemistry, Navjivan Science College, Dahod, 389 151, Gujarat, India

*Corresponding Author E-mail: drktjoshi@gmail.com

,

The condensation reactions of 4-Acetyl/benzoyl-1-(4’-nitrophenyl)-3-methyl-2-pyrazolin-5-ones condensed with sulfamoxole to form bidentate NO donor Schiff bases were studied. The prepared Schiff base ligands were further utilized for the formation of metal chelates having the general formula [ML₂.2H₂O] where M = Mn(II), Fe(II), Co(II), Ni(II), Cu(II) and L = Ligand L₁ and Ligand L₂. These new compounds were characterized by conductance measurements, magnetic susceptibility measurements, elemental analysis, thermo gravimetric analysis, electronic and infrared spectroscopy. Both Schiff base ligands were found to have a mono-anionic bidentate nature and octahedral geometry was assigned to all metal complexes. All the complexes contained coordinated water which was lost at 141–160 °C. These compounds were also screened for their in-vitro antibacterial activity against three bacterial species, namely; Escherichia coli, Bacillus subtilis and Staphylococcus aureus. The metal complexes were found to have greater antibacterial activity than the Schiff base ligands

 

Synthesis of Schiff base ligand:

 

 

 

INTRODUCTION:

4-Aceyl-1-(4’-nitrophenyl)-3-methyl-2-pyrazolin-5-ones belong to a family of heterocyclic b-diketones, and are comparable to b-diketones because in both classes keto-enol tautomerism is possible. Such ligands have played and continue to play a part in the development of coordination compounds that have found a wide application in several fields¹ from new materials to catalysts, as precursors for CVD in the microelectronic industry and as potential antitumourals. The 5-pyrazolone derivatives have been extensively investigated due to their wide range of pharmacological activities².

 

 

 

Sulfonamide derivatives exhibit a range of bioactivities, including anti-angiogenic³, anti-tumor^4,5, anti-inflammatory and anti-analgesic⁶, anti-tubercular⁷, anti-glaucoma⁸, anti-HIV⁹, cytotoxic¹⁰, anti-microbial¹¹ and anti-malarial¹² agents. Sulfonamide derivatives are also known to exhibit a wide variety of pharmacological activities¹³ through exchanges of different functional groups without modification of the structural –S(O)₂N(H)– feature. The synthesis of metal sulfonamide compounds had received much attention due to the fact that sulfanilamides were the first effective chemotherapeutic agents to be employed for the prevention and cure of bacterial infections in humans¹⁴. The pharmacological activity of these types of molecules is often enhanced by complexation with metal ions¹⁵. Moreover, some metal complexes of these ligands have been found to promote rapid healing of burns in humans and animals¹⁶. The effectiveness of burn treatment seems to depend not only on the presence of the metal ion but also crucially on the nature of the material to which the metal ion is bound¹⁷.

 

The coordination behavior of heterocyclic ligands, especially with transition metal ions, has been studied extensively^18-20. Major interest in the complexation of these ligands stems from their suitability as metal-containing model systems, which mimic biologically active system²⁰. The presence of donor atoms (N,S,O) at various positions in these molecules enable them to out as multidentate ligands and thus form chelates of diverse structural types with a wide range of metal ions. The interaction of metal ions with biomolecules and the function of metal ions in physiological systems are very complex, and the precise mechanism of these interactions is almost unknown. In the synthetic systems, ligand design based on selective complexation with metal ions is limited to concepts, such as size-matched selectively in macrocycles, drop in stability due to increased size of chelate and steric strain. However, the ligand activity is a combination of steric, electronic, and pharmakinetic factors, and could be understood in the light of chelation theory. In this context, various heterocycles, especially azoles, occupy an important place owing to their versatile bioactivities due to the presence of multifunctional groups. On the basis of our continuation work^21-26 the present investigation was to synthesize selected Mn(II), Fe(II), Co(II), Ni (II) and Cu(II) metal complexes using sulfamoxole and 4-acetyl/benzoyl-1-(4’-nirophenyl)-3-methyl-2-pyrazolin-5-one as a chelating agent to ascertain their bonding modes and also to study their antibacterial activity.

 

MATERIALS AND METHODS:

Chemicals:

All the reagents used were chemically pure or analytical reagent grade. Solvents were purified and dried according to standard procedures.

 

Synthesis of Schiff base ligands:

The Schiff base ligands were prepared by the condensation of equimolar amounts of 4-acetyl-1-(4’-nitrophenyl)-3-methyl-2-pyrazolin-5-one or 4-benzoyl-1-(4’-nitrophenyl)-3-methyl-2-pyrazolin-5-one and sulfamoxole in minimum quantity of dioxane. The reaction mixture was refluxed in oil bath for three hours. On cooling the crystallized solid Schiff base separate out. It was then filtered and washed with some hot 1,4 dioxane and dried in air. The structure of Schiff base is shown in Scheme 1.

 

Synthesis of metal complexes:

For the preparation of the metal complexes, an aqueous solution of the corresponding metal(II) acetate/sulphate (0.05M) and DMF solution of ligand (0.05M) were mixed in the presence of acetate buffer (pH=6.5) and the mixture was digested on a sand bath at 85-90°C for one hour, cooled and the precipitate filtered and then washed with water and finally with DMF to remove excess metal ion and unreacted Schiff bases.

 

Physical measurements:

Melting points were taken in one side open capillaries on a melting point apparatus VEEGO VMP-D. Electronic Spectra were recorded in DMF solution on a LAMBDA 19, UV/VIS/NIR (“SICART-CVN” at Vallabh Vidyanagar, Gujarat, India). The thermo gravimetric analysis (TGA) was carried out in a dynamic nitrogen atmosphere (20 mL.min⁻¹), with a heating rate of 10 ^◦C min⁻¹ using Shimadzu TGA-50H thermal analyzers. The Mass spectra of all ligands were recorded on a Shimadzu LCMS-2010A. Carbon, Hydrogen and Nitrogen were determined on a Thermo Fisher Flash Elemental Analyzer-1112. The ¹H NMR and ¹³C NMR spectra of all the ligands [in deuterated chloroform (CDCl₃)] were recorded on a BRUKER AVANCE-II 400 spectrometer using TMS [(CH₃)₄Si] as internal standard. The Infrared spectra of the ligands studied in the present work were recorded on a Shimadzu FT-IR-8300 of in KBr (Zydus Research Center, Ahmedabad, India).

 

RESULTS AND DISSCUSSION:

Schiff bases ligand-L₁ (where L₁= 4-acetyl-1-(4’-nitrophenyl)-3-methyl-2-pyrazolin-5-one) and ligand-L₂ (where L₂= 4-benzoyl-1-(4’-nitrophenyl)-3-methyl-2-pyrazolin-5-one) are light yellow colored substance having melting points of 212-214ºC and 228-229°C respectively. The elemental analysis suggests ML₂.2H₂O stoichiometry for all metal complexes. The metal complexes are insoluble in water, but soluble in DMF. All the compounds are solid, colored, stable and non-hygroscopic in nature. All the complexes indicates very low molar conductance, which indicated that the complexes are non-electrolytic in nature. The physical, analytical and conductivity data of the Schiff bases and their metal complexes are shown in Table 1.

Table 1. Physico-analytical data of the ligands and their metal complexes

Compound	Molecular Formula	Molecular Weight	Color	M P (˚C)	λM (Ω-1 cm-1mol-1)	Yield (%)
Ligand-L1	C23H22N6O6S	510.52	Light Yellow	212-214	-	67
Mn(L1)2.2H2O	C46H48MnN12O14S2	1112.01	Brown	~246	16.1	59
Fe(L1)2.2H2O	C46H48FeN12O14S2	1112.92	Radish Brown	~291	12.17	56
Co(L1)2.2H2O	C46H48CoN12O14S2	1116.01	Brown	> 290	13.84	47
Ni(L1)2.2H2O	C46H48NiN12O14S2	1115.77	Green	~276	9.86	49
Cu(L1)2.2H2O	C46H48CuN12O14S2	1120.62	Dark Brown	~261	12.75	56
Ligand-L2	C28H24N6O6S	572.59	Light Yellow	228-229	-	71
Mn(L2)2.2H2O	C56H52MnN12O14S2	1236.15	Brown	~260	11.98	62
Fe(L2)2.2H2O	C56H52FeN12O14S2	1237.06	Radish Brown	~277	8.12	59
Co(L2)2.2H2O	C56H52CoN12O14S2	1240.15	Brown	> 300	9.87	55
Ni(L2)2.2H2O	C56H52NiN12O14S2	1239.91	Green	~286	12.51	57
Cu(L2)2.2H2O	C56H52CuN12O14S2	1244.76	Dark Brown	~248	7.61	48

 

 

FT-IR Spectra:

 From the recorded IR spectra, the wave numbers of the following groups are shown. 3152 cm^-1 (υC-H stretching of Aromatic), 2990 cm^-1 (υC-H stretching of saturated hydrocarbon), 1628 cm^-1, (υC=N stretching of azomethine), 1601 cm^-1, 1495 cm^-1, 1487 cm^-1, 1130 cm^-1 (characteristic bands of pyrazolin ring) and 1507 cm^-1 (Characteristic band of thiazole ring). One of the significant differences to be expected between the IR spectrum of the parent ligand and its metal complex is the presence of more broadened bands in the region of 3500-3100 cm^-1 for the metal complex as the oxygen of the O-H group and nitrogen of the C=N azomethine of the ligand forms a coordination band with the metal ion. The ligand band at 1310 cm^-1 assigned to υ (C-O), shifts to 1340-1315 cm^{-1 27} on complexation, lending further support to the involvement of the nitrogen of the azomethine moiety in the complex formation. The infrared spectra of the ligands show a υO-H (weakly H-bonded) band at 2995 cm^{-1 28}. The absence of this band in all the metal complexes indicates the removal of the proton of the hydroxyl group of the pyrazolin ring during the chelation. The sharp intense band at 1628 cm^-1 in the ligand can be assigned to υC=N (azomethine). A downward shift (∆υ=06-35cm^-1) in υC=N (azomethine) is observed upon coordination, indicating that the nitrogen of the azomethine group is involved in coordination. All the complexes show a broad band in the region 3200 cm^-1 to 3450 cm^-1 which may be assigned to υO-H of coordinated water²⁹. To account for the octahedral stereochemistry of the metal complexes, the coordination of two water molecules is expected. The bands present at ~513 cm^-1 in the Mn(II) complex, ~547 cm^-1 in the Fe(II) complex, ~591 cm^-1 in the Co(II) complex, ~498 cm^-1 in the Ni(II) complex and ~576 cm^-1 in the Cu(II) complex respectively may be due to metal-nitrogen stretching vibrations^30,31. A less intense band at ~1622 cm^-1 in the spectra of the ligands may be assigned to υC=N (ring)³². All the metal complexes do not show shifting in υC=N compared to their respective ligand. This suggests that the nitrogen atom of the thiazole ring has not participated in the coordination. However, in water containing metal complexes, this band is observed as a broad band with some fine structures. This may be due to coupling of the bending mode of coordinated water molecules with υC=N. ³³

 

NMR spectra:

¹H NMR spectrum of ligand L₁: chemical shifts and multiplicities of the corresponding protons are: (400 MHz, CDCl₃) δ = 0.92 (s, 3H, -CH₃), 0.95 (s, 3H, -CH₃), 2.11 (s, 3H, -CH₃), 2.19 (s, 3H, -CH₃), 2.64 (s, 1H), 4.11 (br s, NH), 7.32-8.01 (m, Aromatic Protons). ¹³C NMR Spectrum of ligand L₁: (100 MHz, CDCl₃) δ = 8.1, 10.7, 17.8, 47.9, 121.1, 122.4, 123.3, 124.2, 125.5, 126.8, 138.3, 144.1, 147, 150.7, 152.8, 155.6, 164.9. ¹H NMR Spectrum of ligand L₂: chemical shifts and multiplicities of the corresponding protons are: (400 MHz, CDCl₃) δ = 0.95 (s, 3H, -CH₃), 2.17 (s, 3H, -CH₃), 2.25 (s, 3H, -CH₃), 2.72 (s, 1H), 4.07 (br s, NH), 7.5-8.5 (m, Aromatic Protons). ¹³C NMR Spectrum of ligand L₂: (100 MHz, CDCl₃) δ = 7.8, 17.5, 48.7, 121, 122.5, 123.4, 124, 125.9, 127, 128.2, 129.3, 130.1, 134.1, 144.9, 146.5, 151, 151.9, 155.5, 165.8.

 

Mass spectra:

The positive ion mass spectral analysis of ligand L₁ observed at m/z 511.4 (M⁺), confirms the theoretical molecular weight i.e. 510.52 and for ligand L₂ observed at m/z 573.6 (M⁺), confirms the theoretical molecular weight i.e. 572.59.

 

Electronic spectra:

The electronic absorption spectra are often very helpful in the evaluation of results furnished by other methods of structural investigation. The electronic spectral measurements were used for assigning the stereo chemistries of metal ions in the complexes based on the positions and number of d–d transition peaks. Both the ligands show two absorption bands between 37000 cm^-1 and 26000 cm^-1. No absorption was observed in the visible region for the ligand. In the absence of Quantum mechanical calculation, it is not possible to assign the absorption bands to definite electronic transitions with complete certainty. However, it appears reasonable to assign the bands to π→π* transitions³⁴. The electronic spectrum of the Mn(II) complex exhibits three very low intensity bands, one at 16511 cm^-1, 15208 cm^-1 which may be due to the ⁶A_1g→⁴T_1g (G) transition, another band at 17997 cm^-1, 18164 cm^-1 assigned to the ⁶A_1g→⁴A_1g(G) transition and the third band at 25780 cm^-1, 24554 cm^-1 may be assigned to the  ⁶A_1g→⁴A_1g, 4E_g,(G) transition respectively for the Mn(II) ions in an octahedral environment. The µ_eff (Table 2) value of the complex suggests the 3d⁵ spin configuration³⁵. The electronic spectrum of Fe(II) complex shows broad bands at 23847 cm^-1 and 24419 cm^-1 respectively which may be assigned to the ⁵T_2g→⁵Eg transition. The magnetic moment value for both the ligands 4.97 BM and 4.99 BM respectively indicates that the complex is spin-free and it has octahedral  geometry ³⁶.

 

The electronic spectrum of the Co(II) complexes exhibited two low energy peaks at 8357, 8824 cm^-1 ; 16884, 16948 cm^-1 and a strong high energy peak at 19757, 20016  cm^-1, which can be assigned to the transitions ⁴T_1g(F)→⁴T_2g(F), ⁴T_1g(F)→⁴A_2g(F) and ⁴T_1g(F)→⁴T_2g (P) for a high spin octahedral geometry (Table 3). The magnetic measurements of the Co(II) complexes display in Table 2 magnetic moments is in the octahedral range. The Ni(II) complex exhibited three bands at 10543 cm^-1, 16217 cm^-1 and 25961 cm^-1 for ligand L₁ and 11025 cm^-1, 16161 cm^-1 and 26341 cm^-1 for ligand L₂ which are attributed to the ³A_2g→³T_2g(F) (υ1); ³A_2g→³T_1g(F) (υ2) and ³A_2g→³T_1g(p) (υ3) transitions respectively indicating octahedral geometry around Ni(II) ion. The Ni(II) complex showed the magnetic moment values in the range of 2.90 to 3.00 B.M suggesting consistency with their octahedral environment³⁷. For the Cu(II) complex with D_4h symmetry, three spin allowed transitions ²B_1g→²A_1g (υ1), ²B_1g→²B_2g (υ2)  and  ²B1g→²Eg (υ3) are possible but the electronic spectrum of the Cu(II) complex displayed two bands in the region of 12000 cm^-1 to 14000 cm^-1 and 20000 cm^-1 to 22000 cm^-1. The third transition could not be observed which may be due to very close energy values of the different states. Absence of any spectral bands below 10000 cm^-1 rules out the possibility of a tetrahedral structure of the complexes and also suggests a distorted octahedral geometry of the complexes³⁸.

 

THERMO GRAVIMENTRIC ANALYSIS:

Thermo gravimetric analysis of Schiff base ligand and its metal complexes are used to: (i) Get information about the thermal stability of these new complexes, (ii) Decide whether the water molecules (if present) are inside or outside the inner coordination sphere of the central metal ion and (iii) Suggest the general scheme for thermal decomposition of metal complexes. The data are provided in the Table 4. Number of coordinated or lattice water molecule/molecules present in the complexes were calculated from the percentage weight loss of the complexes from the thermograms. Generally, the loss of lattice water will be at a lower temperature than that of coordinated water^39-42. From the nature of the thermograms and percentage weight-loss, the complexes studied in the present work can be classified in the following three groups^43-45.

 

In the present investigation, heating rates were suitably controlled at 10°C min^-1 under nitrogen atmosphere and the weight loss was measured from the ambient temperature up to ~1000°C^46-48. The thermograms of this group of metal complexes show three stage decomposition. All the metal complexes do not show weight loss below 120°C, and they indicates the absence of lattice water in the metal complexes. The first stage decomposition is obtained in the temperature range 141-160°C. The % weight loss in this range corresponds to the loss of two coordinated water molecules^49-52. The second stage decomposition is obtained in the temperature range 210-400°C. The % weight loss in this range corresponds to % weight loss of two Schiff base ligands. The third stage decomposition range is obtained in the temperature range 400-900°C. The % weight loss in this range corresponds to % weight loss of the metal oxide residue. On the basis of TGA and analytical data, all Mn(II), Fe(II), Co(II), Ni(II) and Cu(II) complexes studied in the present work correspond to the [ML₂.2H₂O] group.

 

ANTIBACTERIAL ACTIVITY:

The Schiff base and their metal complexes were tested for antibacterial activity against Escherichia coli, Bacillus subtilis, Staphylococcus aureus and evaluated by use of the agar disc diffusion method on the basis of the size of the inhibition zone formed around the paper discs. For each concentration, the mean diameter (mm) of the inhibition zone developed was calculated. The test compounds in measured quantities were dissolved in DMF to get concentrations of 200 and 100 ppm of the compounds. Twenty five milliliter nutrient agar media was poured in each Petri dish. After solidification, 0.1mL of test bacteria were spread over the medium using a spreader. The discs of Whatmann no. 1 filter paper, having the diameter 5.00 mm, were placed at four equidistant places at a distance of 2 cm from the center in the inoculated Petri plates. A filter paper disc treated with DMF served as control and Amoxyciline used as a standard drug. These Petri plates were kept in a refrigerator for 24 hours for pre diffusion. Finally, Petri dishes were incubated for 24 hours at 30^oC. The zone of inhibition was calculated in millimeters carefully.

 

The Schiff base ligand was found to be biologically active (Table 5). It is known that chelation tends to make ligands act as more powerful and potent bactericidal agents⁵³. The values indicate that the metal complexes had a higher antibacterial activity than the free ligand. Such increased activity of the metal complexes can be explained on the basis of the overtone concept⁵⁴ and chelation theory⁵⁵. According to the overtone concept of cell permeability, the lipid membrane that surrounds the cell favors the passage of only lipid soluble materials, due to which liposolubility is an important factor controlling the antimicrobial activity. On chelation, the polarity of the metal ion is reduced to a great extent due to the overlap of the ligand orbital and the partial sharing of the positive charge of the metal ion with donor groups. Furthermore, it increases the delocalization of electrons over the whole chelate ring and enhances the lipophilicity of the complex. This increased lipophilicity enhances the penetration of the complex into the lipid membrane and blocks the metal binding sites on the enzymes of the microorganism.

 

 

Table 2. Elemental analysis data of the ligands and their metal complexes

Compound	C % Obs. (Cal.)	H % Obs. (Cal.)	N % Obs. (Cal.)	S % Obs. (Cal.)	M % Obs. (Cal.)	µeff B.M.
Ligand-L1	54.10 (54.11)	4.21 (4.34)	16.16 (16.46)	6.21 (6.28)	-	-
Mn(L1)2.2H2O	49.87 (49.68)	4.22 (4.35)	15.37 (15.11)	5.68 (5.77)	4.51 (4.94)	5.52
Fe(L1)2.2H2O	49.16 (49.64)	4.21 (4.35)	15.08 (15.10)	5.60 (5.76)	4.91(5.02)	4.97
Co(L1)2.2H2O	49.12 (49.51)	4.27 (4.34)	15.07 (15.06)	5.67 (5.75)	4.99 (5.28)	4.48
Ni(L1)2.2H2O	49.47 (49.52)	5.20 (4.34)	15.01 (15.06)	5.70 (5.75)	4.96 (5.26)	2.94
Cu(L1)2.2H2O	49.27 (49.30)	4.30 (4.32)	15.07 (15.00)	5.68 (5.72)	5.07 (5.67)	1.98
Ligand-L2	59.68 (58.73)	4.30 (4.22)	14.58 (14.68)	5.55 (5.60)	-	-
Mn(L2)2.2H2O	54.36 (54.41)	4.21 (4.24)	13.64 (13.60)	5.08 (5.19)	4.12 (4.44)	5.41
Fe(L2)2.2H2O	54.29 (54.37)	4.19 (4.24)	13.51 (13.59)	5.15 (5.18)	4.33 (4.51)	4.99
Co(L2)2.2H2O	54.11 (54.24)	4.20 (4.23)	13.54 (13.55)	5.13 (5.17)	4.67 (4.75)	4.41
Ni(L2)2.2H2O	54.15 (54.25)	4.19 (4.23)	13.22 (13.56)	5.11 (5.17)	4.59 (4.73)	3.00
Cu(L2)2.2H2O	54.10 (54.03)	4.22 (4.21)	13.41 (13.50)	5.13 (5.15)	5.06 (5.11)	1.99

Table 3. Electronic spectroscopic data of the metal complexes

Metals	λmax (cm-1)		Transitions
	M(L1)2.2H2O	M(L2)2.2H2O
Manganese- Mn	16511	15208	6A1g→4T1g (G)
	17997	18164	6A1g→4A1g(G)
	25780	24554	6A1g→4A1g, 4Eg,(G)
Iron- Fe	23847	24419	5T2g→5Eg
Cobalt- Co	8357	8824	4T1g(F)→4T2g(F)
	16884	16948	4T1g(F)→4A2g(F)
	19757	20016	4T1g(F)→4T2g (P)
Nickel- Ni	10543	11025	3A2g→3T2g(F) (υ1)
	16217	16161	3A2g→3T1g(F) (υ2)
	25961	26341	3A2g→3T1g(p) (υ3)
Copper- Cu	12000-14000 20000-22000	12000-14000 20000-22000	2B1g→2A1g (υ1) 2B1g→2B2g (υ2) 2B1g→2Eg (υ3)

 

Table 4. Thermo analytical results of metal complexes

 

Table 5. Antibacterial activity of ligands and complexes

compound	Zone of inhibition in mm(concentration in ppm)
	E. coli		B. subtilis		S. aureus
	100	200	100	200	100	200
Ligand-L1	7	18	8	11	8	11
Mn(L1)2.2H2O	9	12	10	15	10	15
Fe(L1)2.2H2O	10	12	9	16	11	17
Co(L1)2.2H2O	10	17	12	11	11	15
Ni(L1)2.2H2O	17	19	13	20	13	27
Cu(L1)2.2H2O	19	21	15	21	14	24
Ligand-L2	9	14	9	17	9	14
Mn(L2)2.2H2O	9	12	12	15	10	16
Fe(L2)2.2H2O	10	17	11	16	12	17
Co(L2)2.2H2O	10	16	15	15	9	15
Ni(L2)2.2H2O	16	23	21	19	14	24
Cu(L2)2.2H2O	15	22	20	19	13	22
Amoxyciline	18	28	16	22	14	29

 

CONCLUSION:

On the basis of the results obtained from elemental analysis, infrared and electronic spectra, TGA analysis and magnetic susceptibility measurements it is clear that octahedral complex of type [ML₂(H₂O)₂] are formed. Antibacterial activity leads to the following conclusions:

(1)   All metal complexes show more activity than the ligands against the tested bacteria.

(2)   The antibacterial activity of the Cu(II) and Ni(II) complexes have higher activity than the other complexes.

(3)   All metal complexes have an octahedral geometry. 

 

Where, M= Mn, Fe, Co, Ni, Cu

R= -CH₃, -C₆H₅

 

ACKNOWLEDGEMENT:

The authors thank the University of KwaZulu-Natal, School of Chemistry, Durban, South Africa and the Anaj Mahajan Sarvjanik Education Society, Dahod, for their active interest in carrying out this work.

 

REFERENCES:

1.       F. Marchetti, C. Pettinari and R. Pettinari, Coord. Chem. Rev., 249, 2909 (2005).

2.       M. Alaudeen, P.G. Sushama and A.M. Dorothy, Indian J. Chem., A 42, 1617 (2003).

3.       Y. Funahashi, N.H. Sugi, T. Semba, Y. Yamamoto, S. Hamaoka, N. Tsukahara-Tamai, Y. Ozawa, A. Tsuruoka, K. Nara, K. Takahashi, T. Okabe, J. Kamata, T. Owa, N. Ueda, T. Haneda, M. Yonaga, K. Yoshimatsu and T. Wakabayashi, Cancer Res., 62, 6116 (2002).

4.       T. Semba, Y. Funahashi, N. Ono, Y. Yamamoto, N.H. Sugi, M. Asada, K. Yoshimatsu, and T. Wakabayashi, Clin. Cancer Res., 10, 1430 (2004).

5.       J. Sawi´nski and M. Gdaniec, Eur. J. Med. Chem., 40, 377 (2005).

6.       Q. Chen, P.N.P. Rao and E.E. Knaus, Bioorg. Med. Chem., 13, 2459 (2005).

7.       A.K. Gadad, M.N. Noolyi and R.V. Karpoormath, Bioorg. Med. Chem., 12, 651 (2004).

8.       V.K. Agrawal, S. Bano, C.T. Supuran and P.V. Khadikar, Eur. J. Med. Chem. 39, 593 (2004).

9.       C.M. Yeung, L.L. Klein, C.A. Flentge, J.T. Randolph, C. Zhao, M. Sun, T. Dekhtyar, V.S. Stoll and D.J. Kempf, Bioorg. Med. Chem. Lett. 15, 2275 (2005).

10.     I. Enc´ıo, Dj. Morr´e, R. Villar, Mj. Gil and V. Mart´ınez-Merino, Br. J. Cancer, 92, 690 (2005).

11.     M.J. Nieta, F.L. Alovero, R.H. Manzo and M.R. Mazzieri, Eur. J. Med. Chem., 40, 361 (2005).

12.     J.N. Dom´ ınguez, C. Le´on, J. Rodrigues, N.G. de Dom´ ınguez, J. Gut and P.J. Rosethal, IL Farmaco, 60, 307 (2005).

13.     H. Yoshino, N. Ueda, J. Niijima, H. Sugumi, Y. Kotake, N. Koyanagi, K. Yoshimatsu, M. Asada and T. Watanabe, J. Med. Chem.,35, 2496 (1992).

14.     G.G. Mohamed and M.A.M. Gad-Elkareem, Spectrochim. Acta, Part A 68, 1382 (2007).

15.     A. Bult and H. Sigel,In: Metal Ions in Biological Systems, vol. 116. Marcel Dekker, New York, p. 261. (1983).

16.     G.M. De Oliveira, A. Baraldi, L. De Lourenco Marques, E.S. Lang and M.A. Villetti, Inorg. Chim. Acta, 361, 132 (2008).

17.     N.C. Baenziger, S.L. Modak and C.L. Fox Jr, Acta Crystallogr.,Sect. C 39, 1620 (1983).

18.     L. Mishra, A. Jha, H. Itokawa and K. Takeya, Indian J. Chem., 37A, 747 (1998).

19.     Takei, Nishibayashi, Yoshlaki. Chemical Communications, 22, 2360 (2001).

20.     A.Z. El-Sonbati, A.A. El-Bindary, M.A. Diab and S.A. Mazrouh, Monatshefte Fur Chemie. 124, 793 (1993).

21.     K.T. Joshi, A.M. Pancholi, K.S. Pandya, K.K. Singh and A.S. Thakar, Asian J. Chem., 22, 7706 (2010).

22.     A.S. Thakar, K.K. Singh, K.T. Joshi, A.M. Pancholi and K.S. Pandya, E-Journal of Chemistry, 7, 1407 (2010).

23.     K.T. Joshi, A.S. Thakar, A.M. Pancholi and K.S. Pandya, E-Journal of Chemistry, 8, 1556 (2011).

24.     A.S. Thakar, K.T. Joshi, A.M. Pancholi and K.S. Pandya, E-Journal of Chemistry, 8, 1750 (2011).

25.     K.T. Joshi, A.M. Pancholi, K.S. Pandya and A.S. Thakar, J. Chem. Pharm. Res., 3, 741 (2011).

26.     K.T. Joshi, A.M. Pancholi, K.S. Pandya and A.S. Thakar, Int. J. Res. Chem. Environ., 1, 63 (2011).

27.     L.J. Bellamy, The infrared Spectra of Complex Molecules, Chapman and Hall, London p. 108 (1975).

28.     I.A. Patel and B.T. Thaker, Indian J. Chem., 38A, 431 (1999).

29.     A.K. Rana and J.R. Shah, Indian J. Chem., 20A, 615 (1981).

30.     K. Nakamoto, Infrared spectra of Inorganic and coordination Compounds, John Willey, New York (1963).

31.     D.M. Adans, Metal-Ligand and related Vibration, p. 310 Edward Arnold, London (1967).

32.     E.C. Ocafor, J. Inorg. Nucl. Chem., 42, 1155 (1980).

33.     S.M.F. Rahman, J. Ahmad and M.M. Haq, J. Inorg. Nucl. Chem., 35, 1011 (1973).

34.     B.N. Figgis, Introduction to ligand fields, Interscience, New York (1966).

35.     P.P. Dholakiya and M.N. Patel, Synth. React. Inorg. Met.-Org and Nano-met. Chem., 32, 4 (2000).

36.     B.A. Uzoukwu, K. Gloe and H. Duddeck, Indian J. Chem., 37B, 1180 (1998).

37.     R.L. Dutta and A. Shyamal, Element of Magneto Chemistry, 2nd edn. Affiliated East-West press, New Delhi (1982).

38.     F.L. Mani, J. Inorg. Nucl. Chem., 15, 297 (1979).

39.     R.C. Agarwal, N.K. Singh and R.P. Singh, Synth. React. Inorg. Met.-Org. Chem., 14, 637 (1948).

40.     A.C. Fabretti, C.G. Franchini, C.P. Preti and G.Toshi, Can. J. Chem., 55, 344 (1977).

41.     J. Nakagawa and T. Shimanouchi, Spectrochim. Acta, 20, 429 (1964).

42.     D.K. Rastogi and K.C. Sharma, J. Inorg. Nucl. Chem., 36, 2219 (1974).

43.     C.K. Modi, S.H. Patel and M.N. Patel, J. Therm. Anal. Cal., 73, 347 (2003).

44.     P.B. Pansuriya, P. Dhandhukia, V. Thakkar and M.N. Patel, J. Enz. Inhib. Med. Chem., 22, 477 (2007).

45.     N.M. El-Metwally, I.M. Gabr, A.M. Shallaby and A.A. El-Asmy, J. Coord. Chem., 58, 1145 (2005).

46.     N. Deb, S.D. Barudh, N. Sen Savma and N.N. Das, Thermochim. Acta, 53, 320 (1998)

47.     A.A. Soliman, J. Therm. Anal. Cal., 63, 221 (2001).

48.     A.F. Petrovic, D.M. Petrovic, V.M. Zeovac and M. Budimir, J. Thermal. Anal. Cal., 58, 589 (1999).

49.     B.S. Grag, P.K. Singh and S.K. Grag, Synth. React. Inorg. Met.-Org. Chem., 17, 923 (1987).

50.     C.K. Modi, S.H. Patel and M.N. Patel, J. Therm. Anal. Cal., 87, 441 (2007).

51.     G.G. Mohamed and Z.H. Abd El-Wahab, J. Therm. Anal. Cal., 73, 347 (2003).

52.     G.G. Mohamed, F.A. Nour, El-Dien, E.A. Nadia and El-Gamel, J. Therm. Anal. Cal., 67, 135 (2002).

53.     H. Nora and Al-Sha’alan, Molecules, 12, 1080 (2007).

54.     Y. Anjaneyulu and R.P. Rao, Synth. React. Inorg. Met. Org. Chem., 16, 257 (1986).

55.     L. Mishra and V.K. Singh, Indian J. Chem., 32A, 446 (1993).

 

 

Received on 26.04.2013       Modified on 15.05.2013

Accepted on 18.05.2013      © AJRC All right reserved