untitled European Journal of Chemistry 6 (2) (2015) 107‐116 European Journal of Chemistry ISSN 2153‐2249 (Print) / ISSN 2153‐2257 (Online)  2015 Atlanta Publishing House LLC ‐ All rights reserved ‐ Printed in the USA http://dx.doi.org/10.5155/eurjchem.6.2.107‐116.1162 European Journal of Chemistry Journal webpage: www.eurjchem.com Synthesis, spectroscopic, thermal and anti‐microbial studies of transition metal complexes of hydrazone derived from 4,6‐diacetylresorcinol and S‐methyldithiocarbazate Adel Abbas Ahmed Emara 1,2,*, Ali Mahmoud Taha 1, Mahmood Mohamed Mashaly 1 and Omima Mohamed Ibrahim Adly 1 1 Department of Chemistry, Faculty of Education, Ain Shams University, Roxy, Cairo, 11711, Egypt 2 Department of Chemistry, University College in Makkah, Umm Al‐Qura University, Makkah, 21955, Kingdom of Saudi Arabia * Corresponding author at: Department of Chemistry, University College in Makkah, Umm Al‐Qura University, Makkah, 21955, Kingdom of Saudi Arabia Tel.: +966.540.551982. Fax: +966.12.5586531. E‐mail address: adelaaemara@yahoo.com (A.A.A. Emara). ARTICLE INFORMATION ABSTRACT DOI: 10.5155/eurjchem.6.2.107‐116.1162 Received: 26 September 2014 Received in revised form: 30 October 2014 Accepted: 31 October 2014 Published online: 30 June 2015 Printed: 30 June 2015 New series of copper(II), nickel(II), cobalt(II), zinc(II), cadmium(II), iron(III) and oxovanadium(IV) complexes of hydrazone, H3L, ligand derived from the condensation of S‐ methyldithiocarbazate and 4,6‐diacetylresorcinol, in the molar ratio 1:1, has been synthesized. All the metal complexes are dimmers. The structures of the ligand and its transition metal complexes were characterized by elemental analyses, spectral (Infrared, electronic, Mass, 1H NMR and ESR) data and magnetic susceptibility, molar conductivity measurements and thermal gravimetric analysis. The structure of the ligand is dibasic tridentate with ONS sites. The bonding sites, in all cases, are the azomethine nitrogen, phenolic oxygen and thiol sulfur atoms, as illustrated from the spectral data. The metal complexes exhibit different geometrical arrangements such as square planar, tetrahedral, square pyramidal and octahedral arrangements. Kinetic parameters (G, H, S and E) of the metal complexes were calculated from the thermal behaviour of the metal complexes using Coats‐Redfern method. The ligand and its metal complexes were screened for its antimicrobial activity against Staphylococcus aureus and Staphylococcus pyogenes as Gram‐positive bacteria, Pseudomonas phaseolicola and Pseudomonas fluorescens as Gram‐negative bacteria and the fungi Fusarium oxysporum and Aspergillus fumigatus. KEYWORDS Tridentate ligand Kinetic parameters 4,6‐Diacetylresorcinol Coats‐Redfern method Antimicrobial activities S‐Methyldithiocarbazate Cite this: Eur. J. Chem. 2015, 6(2), 107‐116 1. Introduction Thiosemicarbazones of S‐alkyldithiocarbazate are one of the most important classes of ligands, which have considerable pharmacological interest due to significant biological activities [1‐6]. These compounds have received much attention and interest for further study because they provide an interesting series of ligands whose properties can be greatly modified by introducing different organic substituents, thereby causing a variation in the ultimate donor properties [7]. The presence of hard nitrogen and soft sulfur donor atoms in the backbones of these ligands enable them to react readily with transition metal ions yielding colored stable metal complexes [8‐12]. Moreover, these complexes are potentially biologically active and exhibit antifungal, antimicrobial, antimalarial, antitumoral, antiviral, and antioxidant activities [13‐20]. In this manuscript, condensation of S‐methyldithio carbazate with 4,6‐diacetylresorcinol, in the molar ratio 1:1, afforded the corresponding S‐methyl‐2‐[1‐(5‐acetyl‐2,4‐di hydroxyphenyl)ethylidene]hydrazine carbodithiozate, H3L, ligand. The reactions of H3L ligand with transition metal ions in the molar ratio (ligand:metal ion; 1:1) were studied. The newly prepared metal complexes of this ligand were identified by different physicochemical and spectroscopic techniques. Kinetic parameters (G, H, S and E) of the metal complexes were calculated from the thermal behavior of the metal complexes using Coats‐Redfern method. The ligand and its metal complexes were tested against Staphylococcus aureus (ATCC 25923) and Staphylococcus pyogenes (ATCC 19615) as Gram‐positive bacteria, Pseudo‐ monas phaseolicola (GSPB 2828) and Pseudomonas fluorescens (S 97) as Gram‐negative bacteria and the fungi Fusarium oxysporum and Aspergillus fumigatus. 2. Experimental 2.1. Materials S‐methyldithiocarbazate [21] and 4,6‐diacetylresorcinol [22,23] were prepared as cited in the literature. Copper(II), nickel(II), cobalt(II), zinc(II), cadmium(II) and iron(III) were used as nitrate salts and were Merck or BDH. 108 Emara et al. / European Journal of Chemistry 6 (2) (2015) 107‐116 Scheme 1 Oxovanadium(IV) sulphate monohydrate was BDH. Lithium hydroxide monohydrate was BDH. Organic solvents [Ethanol, methanol, diethylether, acetone, dimethylformamide (DMF) and dimethylsulfoxide (DMSO)] were reagent grade. 2.2. Synthesis of H3L ligand The H3L ligand was synthesized by the condensation of S‐ methyldithiocarbazate (1.22 g, 10.0 mmol) dissolved in hot absolute ethanol (30 mL) with 4,6‐diacetylresorcinol (1.94 g, 10.0 mmol) in absolute ethanol (30 mL). The reaction mixture was heated to reflux for 30 min. The yellow crystals were obtained after cooling the reaction mixture to room tempe‐ rature. The product was filtered off and washed with few portions of ethanol then diethylether and air‐dried. Fine crystals were obtained by recrystallization from ethanol. The ligand was kept in a desiccator until used. The yield was 1.75 g (59 %) and the melting point (M.p.) was 226‐227 C. Scheme 1 illustrates the formation of the H3L ligand. 2.3. Synthesis of the metal complexes Methanolic solutions of the metal salts (30 mL) were added gradually to methanolic solutions of the deprotonated ligand (30 mL) in the molar ratio 1:1. In each case, the ligand was deprotonated using lithium hydroxide, LiOH.H2O, as deprotonating agent. The reaction mixture of the deproto‐ nated, H3L, ligand with metal salts was heated to reflux for one to three hours. The time of the reflux depends on the formation of the solid products. The resulting precipitates were filtered off, washed with methanol then ether. Lithium hydroxide, LiOH.H2O (0.378 g, 9.0 mmol) dissolved in methanol (20 mL) was added to the H3L ligand (0.894 g, 3.0 mmol) in methanol (30 mL), i.e., in the molar ratio 3:1 (LiOH.H2O:ligand) and heated to reflux for 30 min. The appropriate weight (3.0 mmol) of the metal salts was added to the ionic deprotonated ligand to form the metal complexes. The following detailed preparations are given as examples and the other complexes were obtained similarly. 2.3.1. Reaction of nickel(II) with H3L ligand Nickel(II) nitrate hexahydrate, Ni(NO3)2.6H2O, (0.870 g, 3.0 mmol) in methanol (30 mL) was added gradually with constant stirring to a solution of the deprotonated, H3L, ligand (0.894 g, 3.0 mmol) in methanol (30 mL). The stoichiometry of the metal ion to ligand was 1:1. The solution was heated to reflux for three hours. A red precipitate was formed on cold and washed with small portions of methanol then ether. The yield was 1.42 g (69%) and the melting point was over 320 C. 2.3.2. Reaction of oxovandium(IV) with H3L ligand Oxovandium(IV) sulphate monohydrate, VOSO4.H2O, (0.489 g, 3 mmol) in a least amount of distilled water, just to dissolve VOSO4.H2O, was added gradually with constant stirring to a solution of the deprotonated H3L ligand (0.894 g, 3 mmol) in methanol (60 mL). The solution was heated to reflux for three hours. A green precipitate was obtained and washed with small portions of methanol then ether. The yield was 1.56 g (69.03%) and the melting point was over 320 C. 2.4. Physical measurements The analyses of carbon, hydrogen, nitrogen and sulfur were carried out at the National Research Center, Giza, Egypt. Analyses of metal ions were performed after the dissolution of the solid complex in hot concentrated nitric acid, HNO3. The metal ions were then titrated with EDTA [24‐26]. The FT‐IR spectra (4000‐400 cm‐1) of the compounds were recorded as KBr discs using FT‐IR (Shimadzu) spectrophotometer model 4000. Absorption frequencies are given in wavenumbers (cm‐1). 1H NMR spectra were recorded using a Varian spectro‐ meter, 200 MHz. DMSO‐d6 was used as a solvent and tetramethylsilane (TMS) as an internal reference. The spectra were extended from 0‐15 ppm. The chemical shifts (δ) are given down field relative to TMS. D2O was added to every sample to test for the deuteration of the samples. The electronic spectra of the ligand and their metal complexes were carried out on a JASCO model V‐550 UV‐VIS spectro‐ photometer in the range 200‐900 nm. Magnetic susceptibilities of the complexes were measured by the Gouy method at room temperature using Johnson Matthey, Alfa product, Model No. (MKI), Magnetic Susceptibility Balance. The diamagnetic corrections were calculated from Pascal’s constants for all atoms in the compounds [27]. Conductivities were measured in solutions of the complexes in DMF (1×10‐3 M) using WTWD‐ 812 Weilheium‐conductivity meter model LBR, fitted with a cell model LTA100. Thermal gravimetric analysis (TGA) data was measured from room temperature up to 800 C at a heating rate of 20 C/min. The data were obtained using a Shimadzu TGA‐50H instrument. All melting points were recorded in open capillary tubes on a Stuart SMP3 melting point apparatus. ESR spectra of compounds were recorded on the Bruker, Model: EMX, X‐band spectrometer. Mass spectra of the compounds were recorded on a Hewlett Packard mass spectrometer model MS 5988. Samples were introduced directly to the probe, and the fragmentations were carried out at 300 °C and 70 eV. 2.5. Biological activity The standardized disc‐agar diffusion method [28] was followed to investigate the activity of the current compounds against the sensitive organisms Staphylococcus aureus (ATCC 25923) and Streptococcus pyogenes (ATCC 19615) as Gram‐ positive bacteria, Pseudomonas phaseolicola (GSPB 2828) and Pseudomonas fluorescens (S 97) as Gram‐negative bacteria and the fungi Fusarium oxysporum and Aspergillus fumigatus. The antibiotic cephalothin was used as standard reference control in the case of Gram‐positive bacteria and chloramphencol was used as standard reference control in the case of Gram‐ negative bacteria, cycloheximide was used as standard antifungal reference control. The tested compounds were dissolved in DMF to get concentrations of 1 and 2 mg/mL. Emara et al. / European Journal of Chemistry 6 (2) (2015) 107‐116 109 Table 1. Physical and analytical data of the hydrazone, H3L, ligand and its transition metal complexes. Ligand/Complex Molecular formula M.wt. (Yield %) Color M.p. (°C) Elemental Analyses, Found (Calcd.) % C H N S M H3L C12H14N2O3S2 298.37 (59) Yellow 226‐227 47.83 (48.30) 4.70 (4.72) 9.36 (9.39) 22.01 (21.49) ‐ [Cu(HL)]2 C24H24N4O6S4Cu2 719.72 (65) Dark Green >320 39.61 (40.00) 3.56 (3.35) 7.60 (7.89) 18.20 (17.82) 17.52 (17.64) [Ni(HL)]2.2H2O C24H28N4O8S4Ni2 746.16 (63) Red >320 37.83 (38.63) 3.52 (3.78) 7.59 (7.50) 17.04 (17.18) 15.69 (15.73) [Co(H2L)(H2O)2]2(NO3)2.2H2O C24H38N6O18S4Co2 946.72 (70) Green >320 30.20 (30.44) 3.92 (4.04) 8.65 (8.87) 13.20 (13.54) 12.80 (12.65) [VO(HL)]2 C24H24N4O8S4V2 726.50 (69) Green >320 39.70 (39.67) 3.65 (3.30) 7.76 (7.70) 18.00 (17.61) NA* (14.01) [Fe(H2L)(CH3OH)NO3]2.2H2O.CH3OH C27H40N6O17S4Fe2 960.87 (57) Red >320 34.28 (33.74) 3.96 (4.19) 8.25 (8.74) 12.41 (13.34) 11.30 (11.65) [Zn(HL)]2.6H2O C24H36N4O12S4Zn2 830.81 (69) Yellow >320 34.05 (34.69) 4.82 (4.36) 6.24 (6.74) 14.32 (15.43) 15.16 (15.64) [Cd(HL)]2.H2O C24H26N4O7S4Cd2 834.74 (76) Yellow >320 35.01 (34.50) 3.30 (3.11) 5.97 (6.71) 15.03 (15.36) 26.52 (26.83) Scheme 2 The test was performed on medium potato dextrose agar (PDA) which contains infusion of 200 g potatoes, 6.00 g dextrose and 15.0 g agar [29]. Uniform size filter paper disks (3 disks per compound) were impregnated by equal volume (10 µL) from the specific concentration of dissolved tested compounds and carefully placed on incubated agar surface. After incubation for 36 h at 27 °C in the case of bacteria and for 48 h at 24 °C in the case of fungi, inhibition of the organisms which evidenced by clear zone surround each disk was measured and used to calculate mean of inhibition zones [28,29]. The activity of tested compounds was categorized as follows: (i) Low activity = Mean of zone diameter ≤ 1/3 of mean zone diameter of control. (ii) Intermediate activity = Mean of zone diameter ≤ 2/3 of mean zone diameter of control. (iii) High activity = Mean of zone diameter > 2/3 of mean zone diameter of control. 3. Results and discussion 3.1. The hydrazone H3L ligand The hydrazone, H3L, ligand was synthesized by the condensation of S‐methyldithiocarbazate with 4,6‐diacetyl resorcinol in the molar ratio 2:1 or 1:1. In both cases, the obtained product was found to be 1:1 as investigated from elemental analysis and spectral data. The structure of the hydrazone, H3L, ligand was identified by elemental analyses, infrared, UV‐Visible, 1H NMR and mass spectra. Table 1 lists the physical and analytical data of the H3L ligand and its transition metal complexes. It is expected that the H3L ligand has two tautomers, thione and thiol, forms. The thione is the dominant form in either solid state or solution, while the thiol form became more dominant and mainly the most effective in presence of metal ions [30‐32]. The H3L ligand is dibasic tridentate with ONS sites. Scheme 2 is representative struc‐ tures of the thione‐thiol forms of the H3L ligand. The infrared frequencies of the present H3L ligand along with S‐methyldithiocarbazate (SMDTC), 4,6‐diacetylresorcinol (DAR) and their tentative assignments are listed in Table 2. The infrared spectra are consistent with the formation of H3L ligand. The vibrational assignments were aided by comparison with the vibrational frequencies of the related compounds, such as, the hydrazones of S‐methyldithiocarbazate and hydrazones of 4,6‐diacetylresorcinol [11,33‐36]. The funda‐ mental stretching mode of the azomethine moiety, ν(C=N), is readily assigned by comparison with the infrared spectra of S‐ methyldithiocarbazate and 4,6‐diacetylresorcinol. This intense band at 1602 cm‐1 for H3L ligand was assigned to the C=N stretching frequency of the ligand and are characterized for the azomethine moiety of most hydrazone compounds. The absorption band of the C=O appeared at 1643 cm‐1 in the infrared spectrum of the ligand, which gave an indication that one of the C=O groups in the 4,6‐diacetylresorcinol was involved in the condensation reaction while the other one still free, and its stretching frequency still persist with slightly lower frequency than the free 4,6‐diacetylresorcinol, Δν = 24 cm‐1. The ν(NH2) stretching frequencies of S‐methyldithio carbazate which lies at 3307 and 3241 cm‐1 were also disappeared after the formation of the H3L ligand. One band was observed at 3319 cm‐1 which is assigned to ν(NH) in the infrared spectrum of the ligand. On the other hand, other fundamental bands were assigned in the infrared spectra of H3L ligand which are listed in Table 2. The mass spectrum of the hydrazone, H3L, ligand, revealed the molecular ion peaks at m/e 298, which is coincident with the formula weight (298.37) for H3L ligand and supports the identity of the structure. The fragmentation pattern of the mass spectrum of the H3L ligand is depicted in Scheme 3. 1H NMR spectra of H3L ligand in deuterated dimethyl sulphoxide, DMSO‐d6, without and with D2O are shown in Figure 1. The chemical shifts of the proton signals in the spectrum of the hydrazone, H3L, ligand with their assignments are listed in Table 3. Both signals of the ‐OH for the phenolic protons which lies at 12.56 ppm and signal of the ‐NH proton which lies at 12.36 ppm in the ligand were disappeared in the presence of D2O. 1H NMR spectra do not show any resonance near to ~4.0 ppm attributable to the S‐H proton resonance, indicating that the thione form still persists in solution. Electronic spectral data of the H3L ligand was recorded in DMF solution. Two absorption bands were observed at 299 and 355 nm. The irst band corresponds to the π → π* transition of the C=N, C=S and C=O groups, and the second band corresponds to the n → π* transitions resulting from nitrogen, oxygen and sulfur atoms, which are overlapped with the intermolecular charge transfer (CT) from the phenyl ring [37‐39]. 110 Emara et al. / European Journal of Chemistry 6 (2) (2015) 107‐116 Table 2. Infrared frequencies (cm‐1) of H3L ligand, S‐methyldithiocarbazate (SMDTC) and 4,6‐Diacetylresorcinol (DAR), and their tentative assignments *. H3L SMDTC DAR Assignments 3423 w, br ‐ 3222 w, br ν(OH) ‐ 3307 m, 3241 m ‐ ν(NH2) 3319 s 3179 m, br ‐ ν(NH) 3043 w ‐ 3083 w ν(ArC‐H) 2921 m 2914 w - ‐ 1643 s - 1667 vs ν(C=O) 1602 w - - ν(‐C=N‐) 1469 m, 1421 w 1469 s 1490 m δ(CH3) asym. and δ (CH2) asym. 1338 m 1340 m - ν(C=S) 1262 w - 1258 vs δ(OH) in plane 1185 w 1170 w 1180 m ν(C‐N) and ν (C‐C) 1068 s 1046 s - ν(N‐N) 1028 w - 1030 w, sh δ(ArC‐H) in plane 952 m, 890 w 966 w, 910 w 951 w, 903 w ρr(CH3) and ρr(CH2) 510 w ‐ 423 w δ(OH) out of plane * s = strong, m = medium, w = weak, vs = very strong, sh = shoulder, br = broad. OHHO N H N SCH3 S CH3 O CH3 OHHO N H N SCH3 S CH3 O + OHHO N H N S CH3 O CH3 + OHHO N CH3 O CH3 + N H N SCH3 S CH3 OHHO O CH3 + + HC N H N SCH3 S + OH + + + OHHO NHO CH3 + OHHO N O CH3 + OHHO N N C S CH3 O CH3 OHHO N N C S CH3 OHHO O CH3 OHHO d d a a b b c c - . SCH3 - 47 - .CH3 - 15 M 298 [298; 2%] M + 1 299 [299; 1%] - HSCN - 59 251 [251; 5%] 192 [192; 24%] - 15 + 1 - .CH3 + H. - 1 - H. 250 [250; 17%] 178 [178; 100%] 147 [147; 11%]151 [151; 10%]152 [152; 12%] + H. + 1 283 [283; 6%] -1 - H.- .COCH3- 43 - CO - 28 - .COCH3 - H. - 43 + 1 - .OH - .COCH3- 43 + 1 - 17 - CH3- 15 177 [177; 13%] 207 [207; 10%] 133 [133; 12%]93 [93; 16%]65 [65; 8%]110 [110; 12%] + H. + H.+ 1 Scheme 3 Figure 1. 1H NMR spectra (, ppm) in DMSO‐d6 solvent of: (a) H3L ligand without addition of D2O, (b) H3L ligand after the addition of D2O and (c) [Zn(HL)]2.6H2O complex (6); (*) Suppressed H2O. 3.2. Metal complexes of H3L ligand The hydrazone, H3L, ligand behaves as a dibasic with tridentate ONS sites. The ligand reacts with Cu(II), Ni(II), Co(II), Zn(II), Cd(II), Fe(III) and VO(IV) ions to yield the corresponding transition metal complexes. The isolated tran‐ sition metal complexes were identified by elemental analyses, FT‐IR, 1H NMR, mass, UV‐Visible, ESR spectroscopy and thermal gravimetric analysis (TGA), beside, the magnetic susceptibility and molar conductivity measurements. The physical and analytical data of the transition metal complexes are listed in Table 1. 3.2.1. Infrared spectra The characteristic vibrational frequencies and their assignments for the synthesized transition metal complexes are listed in Table 4. The assignments were aided by comparison with the vibrational frequencies of the free ligand and other related compounds such as transition metal complexes of 4,6‐diacetylresorcinol hydrazones and 6‐methyl‐ 4‐formylpyrimidine of S‐methyldithiocarbazate [33,40]. Ins‐ pection of the data reveals that, the stretching frequencies of azomethine group of the metal complexes, ν(‐C=N‐), are shifted to lower frequencies at the range 1563‐1535 cm‐1 and are in lower frequencies, than the corresponding ν(‐C=N‐) frequency of the H3L ligand which lie at 1602 cm‐1. This lowering shift may be due to the coordination of the azomethine group to the metal ion [41‐43]. Emara et al. / European Journal of Chemistry 6 (2) (2015) 107‐116 111 Table 3. The 1H NMR chemical shifts (ppm) of the hydrazone, H3L, ligand and its [ZnHL]2.6H2O complex (6) in DMSO‐d6 and their assignments *. HO N N SCH3 SO Zn CH3CH3 O OH N N SCH3 S O CH3 CH3 O (a) (g) (c) (d) (e) (b) 6H2O Zn H3L [ZnHL]2.6H2O complex (6) Chemical shifts (ppm) Assignments H3L [ZnHL]2.6H2O (6) 2.52 2.48 (s, 3H, SCH3) (a) H3L, (s, 6H, 2 SCH3) (a) (6) 2.61 2.59 (s, 3H, CH3) (b) H3L, (s, 6H, 2 CH3) (b) (6) 2.64 2.69 (s, 3H, CH3) (c) H3L, (s, 6H, 2 CH3) (c) (6) 6.39 6.04 (s, 1H, 1 Ar‐H) (d) H3L, (s, 2H, 2 Ar‐H) (d) (6) 8.13 8.12 (s, 1H, 1 Ar‐H) (e) H3L, (s, 2H, 2 Ar‐H) (e) (6) 12.36 ‐ (s, 1H, NH) exchangeable by D2O (f) H3L 12.56 12.67 (s, 1H, OH) exchangeable by D2O (g) H3L, (s, 2H, 2 OH) exchangeable by D2O (g) (6) * s = singlet, Data given for the spectrum depicted in Figure 1. Chemical shifts (ppm) were referenced at 25 °C with respect to TMS. Table 4. Characteristic infrared frequencies (cm‐1) a of the hydrazone, H3L, ligand and its transition metal complexes. Ligand/ Complex ν(OH) ν(C=O) ν(C=N) ν(CH) aliph. ν(CH) Ar. ν(C‐S) ν(C‐S‐C) ν(N‐N) ν(M‐N) ν(M‐O) Other significant bands H3L b 3423 w, br 1643 s 1602 s 2921m, 2851 w 3043 w ‐ 952 m 1068 m ‐ ‐ ‐ [Cu(HL)]2 3447 w, br 1640 s 1540 s 2920 w 3093 w 676 w 993 m 1085 w 604 m 485 w ‐ [Ni(HL)]2.2H2O 3440 w, br 1638 s 1535 s 2920 m ‐ 687 w 958 w 1084 w 607 m 467 w ‐ [Co(H2L)(H2O)2]2(NO3)2.2H2O 3423 s, br 1629 m 1539 s 2927 w ‐ 670 w 964 w 1085 w 602 w 470 w NO3‐ (ionic) 1754 (s), 1384 (s), 836 (s) [VO(HL)]2 3404 w, br 1642 s 1538 s 2918 w 3091 w 677 w 957 m 1083 w 604 w 485 w 993 (s) for ν(V=O) [Fe(H2L)(CH3OH)NO3]2.2H2O. CH3OH 3386 s, br 1642 s 1563 s 2927 w ‐ 665 w 962 m 1085 w 576 w 491 w NO3‐ (unidentate) 1426, 1320, 809 [Zn(HL)]2.6H2O 3422 s, br 1633 s 1554 s 2927 w ‐ 679 w 964 m 1085 w 590 m 488 w ‐ [Cd(HL)]2.H2O 3420 s, br 1618 s 1548 s 2926 w ‐ 670 w 966 m 1084 w 580 m 470 w ‐ a s = strong, m = medium, w = weak, br = broad. b Stretching frequencies of the ν(NH) and ν(C=S) of the H3L ligand which observed at 3319 and 1338 cm‐1; respectively, were disappeared in all the infrared spectra of the metal complexes. The broad bands in the range of 3447‐3386 cm‐1 can be assigned to the stretching frequencies of the ν(OH) of water and/or methanol molecules associated to the complexes which are also confirmed by the elemental analyses. The ν(C=S) band at 1338 cm‐1 for H3L ligand was not observed in the metal complexes, thus supporting the suggestion of coordination through the thiol sulfur and showed the ν(C‐S) band in the range of 687‐665 cm‐1. The shift frequency of ν(N‐N) band in all the metal complexes to higher frequency in the range 1085‐ 1083 cm‐1 compared to the ν(N‐N) band in the free ligand at 1068 cm‐1, is another evidence to that the coordination involves thiol sulfur and not the thione sulfur. The weak bands in the ranges 491‐467 cm‐1 and 607‐576 cm‐1 were assigned to the ν(M‐O) and ν(M‐N) stretching frequencies of the metal‐ oxygen and metal‐nitrogen bonds for the transition metal complexes of the H3L ligand [44]. This emphasized supporting that the bonding sites of the ligand to the metal ions is achieved by the phenolic oxygen atom, the azomethine nitrogen atom and thiol sulfur atom of the ligand. The free NO3– exhibited three bands at 1754, 1381 and 836 cm‐1 in [Co(H2L)(H2O)2]2(NO3)22H2O complex (3), and the unidentate NO3– ions exhibited three bands at 1426, 1320 and 809 cm‐1 in [Fe(HL)(CH3OH)(NO3)]2.H2O.CH3OH complex (5) [44,45]. The IR spectrum of VO(IV) complex (4) displayed a band at 993 cm‐1 which have no counterpart in the spectrum of the H3L ligand and is assigned to the stretching frequency of the ν(V=O) band [44]. 3.2.2. Electronic, ESR, mass spectra and magnetic moment measurements and molar conductance Table 5 lists the electronic spectral bands of the metal complexes in Nujol mulls and in DMF solution and the magnetic moments of the metal ions in their complexes. The absorption spectra exhibit no obvious difference on going from Nujol mulls to DMF solution. The electronic spectrum of the dimer dark green [Cu(HL)]2 complex (1) showed one band due to Eg → 3T2g(G) transition, which was observed at 598 nm which confirms square planar structure [46,47]. The subnormal magnetic moment of Cu(II) complex representing in a strong metal‐metal interaction. Elemental analysis and spectral data confirm square planar structure for the Cu(II) complex [46‐48]. X‐Band ESR spectra of [Cu(HL)]2 (1) complexes are silent due to strong metal‐ metal bond in dimeric form. The electronic spectrum of the red [Ni(HL)]2.2H2O complex (2), is consistent with the square planar transitions. One electronic transition band was observed at 499 nm, which was assigned to 1A1g → 1A2g transition [49]. In most cases, it is rarely possible to identify the d‐d transitions of nickel(II) square planar complexes in presence of organic ligands, since even the very weak tail of UV‐organic absorption tailing into the visible can obscure them. 112 Emara et al. / European Journal of Chemistry 6 (2) (2015) 107‐116 Table 5. Electronic (d‐d) transition bands (nm) a and magnetic moments (B.M.) of the transition metal complexes of the hydrazone, H3L, ligand. Complex λ max (nm) Nujol mulls Assignments Magnetic moments (B.M.) Conductance Geometry (DMF nm; ε max) μcompl. b μeff. c ohm‐1·cm2·mol‐1 [Cu(HL)]2 598 (589; 0.0058) Eg → 3T2g(G) d Diamag. Diamag. 1.35 Square planar [Ni(HL)]2.2H2O 499 (423; 0.0036) 1A1g → 1A2g Diamag. Diamag. 2.65 Square planar [Co(H2L)(H2O)2]2(NO3)22H2O 454 (420; 0.0070), 608 (600; 0.0043) 4T1g(F) → 4T1g(P) e 4T1g(F) → 4A2g(F) 5.98 4.70 141.00 Octahedral [VO(HL)]2 591 (595; 0.0048) b2 → b1 Diamag. Diamag. 1.65 Square pyramid [Fe(H2L)(CH3OH)NO3]2.2H2O.CH3OH 518 (501; 0.0067), 553 (523; 0.0088) Charge transfer (CT) band tailing from UV to visible region. 3.73 2.64 14.00 Octahedral [Zn(HL)]2.6H2O ‐ ‐ Diamag. Diamag. 2.65 Tetrahedral [Cd(HL)]2.H2O ‐ ‐ Diamag. Diamag. 4.52 Tetrahedral a Electronic (d‐d) transition bands of the complexes are in (nm) and were recorded in the solid state as Nujol mulls and DMF value in parentheses are the wave number in DMF solvent and ε max values were multiplied by 10‐4 (mol‐1·cm‐1). The electronic absorption bands due to the hydrazone ligand in the complexes were observed in each spectrum in the UV region and were not cited in the table. b μcompl. is the magnetic moment of all metal ions present in the complex. c μeff. is the magnetic moment of one metal ion in the complex. d Racha parameter (Dq) of Cu(II) (1) complex is 589 nm. e Racha parameters for Co(II) complex (3) , B, Dq and 1 are 1.02, 1144, 915 and 1266; respectively. The red nickel(II) complex (2) is diamagnetic suggesting a square planar geometry around the nickel(II) ion. The electronic spectrum of the green Co(II) complex (3) showed that, the complex has two bands at 454 and 608 nm indicating an octahedral geometry. The transitions can be interpreted by using Tanab‐Sugano diagram. In the Co(II) octahedral complexes, the spectra usually consist of three bands [50]. The third band, which is due to 4T1g(F) → 4T2g(F) transition, was not observed due to the fact that it occurs in the near infrared region, and out of the range of the used instrument. The other two bands were observed in the visible region, The former band due to 4T1g(F) → 4T1g(P) transition was observed at 454 nm and is slightly obscured by UV‐ organic ligand transitions. The second one due to 4T1g(F) → 4A2g(F) transition, was observed at 608 nm for complex (3). Rach parameters for complexes (3) were calculated which gave the values 1.02, 1144, 915, 1266 for , B, Dq, 1; respectively. The Co(II) complex showed the magnetic moment 4.7 B.M., emphasizing that the complex is octahedral [51]. The electronic spectrum of the green oxovanadium(IV) complex (4) showed one band at 591 nm; which corresponds to b2 → b1 electronic transition [52]. With the aid of the elemental analysis and the infrared spectra, the oxovanadium(IV) complex (4) is 5‐coordinate which would be square pyramid [53]. Oxovanadium(IV) complex is diamagnetic due to strong V‐V bond in dimeric form. X‐Band ESR spectra of [VO(HL)]2 (4) complexes are silent due to strong metal‐metal bond in the dimeric form. The electronic spectrum of the red Fe(III) complex (5) showed two medium bands at 518 and 553 nm. It was not possible to identify the type of the d‐d transition, due to a strong charge transfer (CT) band tailing from the UV‐region to the visible region [51]. Generally, from the elemental and infrared spectrum which gave significant proof for the nitrate anion to act as a unidentate ligand, it is expected that the Fe(III) complex (5) has octahedral arrangement. Magnetic moment of the Fe(III) complex was measured and gives 5.34 B.M. This value refer to that the Fe(III) complex (5) is high spin octahedral geometry [51]. The yellow [Zn(HL)]2.6H2O (6) and [Cd(HL)]2.H2O (7) complexes, are diamagnetic as expected and their geometries are most probably tetrahedral for Zn(II) complex (6) and Cd(II) complex (7). 1H NMR spectrum of [Zn(HL)]2.6H2O (6) in DMSO‐d6 is depicted in Figure 1. The chemical shifts of the proton signals in the spectrum and their assignments are listed in Table 3. It is worthnoting that the signal of one phenolic proton was observed in the spectrum while the other phenolic and amino (thiol tautomer) protons were not observed in the 1H NMR spectrum of the Zn(II) complex, which means that these protons were deprotonated to coordinate with the Zn(II) ion. The molar conductance of the complexes was measured in DMF, 1×10‐3 M, which gave low values (12.52‐1.35) ohm‐ 1cm2mol‐1. However, the molar conductance value Co(II) complex (3) are 141 ohm‐1cm2mol‐1 which fall within the range of 1:1 electrolytes which indicate that NO3‐ ion is completely in the outer sphere [54]. 3.2.3. Thermal gravimetric analysis Figure 2 depicts the TGA‐DrTGA curves of (A) [Fe(HL)(CH3OH)(NO3)]2.H2O.CH3OH complex (5) and (B) [Cd(HL)]2.H2O complex (7). Thermal gravimetric analysis (TGA) for [Ni(HL)]2.2H2O complex (2), shows four stages. The first stage is due to the loss of two uncoordinated water molecules (weight loss; Calcd./Found%; 4.82/4.83%) which observed at 27‐100 C. The second stage was observed at 212‐ 324 C and is due to the loss of (CH3S)2 molecule. (weight loss; Calcd./Found%; 12.59/ 12.35%). The third stage was observed at 324‐523 C and is due to loss of one (CN)2 molecule of two CH3CN molecules and CH3CHO from the complex molecules (weight loss; Calcd./Found%; 29.48/ 28.76%). The fourth stage was observed at 523‐800 C and is due to loss C6H8O2 molecules (weight loss; Calcd./Found%; 14.74/14.42%). The residue due to the decomposition of the rest part of the complex was [NiOSC2]2 (weight loss; Calcd./Found%; 38.24/38.90%). TGA of [Co(H2L)(H2O)2]2(NO3)2.2H2O complex (3), shows three stages. The first stage is due to the loss of two uncoordinated water molecule (weight loss; Calcd./Found%; 3.80/3.88%) which observed at 26‐120 C. The second stage was observed at 175‐329 C and is due to the loss of four coordinated water molecules and two CH3SCN molecules (weight loss; Calcd./Found%; 23.02/23.49%). The third stage was observed at 330‐800 C and is due to loss of two CH3CN, N2O5, CH3CO and C6H3OH molecules (weight loss; Calcd./ Found%; 40.56/41.13%). The residue due to the decompo‐ sition of the rest part of the complex showed some organic fragment associated with the metal oxide and metal sulfide (weight loss; Calcd./Found%; 32.30/32.75%). TGA of [Fe(HL)(CH3OH)(NO3)]2.H2O.CH3OH complex (5), shows three stages. Figure 2A depicts the TGA‐DrTGA curves of [Fe(HLc)(CH3OH)(NO3)]2.H2O.CH3OH complex (5). The first stage is due to the loss of one uncoordinated methanol molecule and one uncoordinated water molecule (weight loss; Calcd./Found%; 5.31/5.00%) which observed at 52‐100 C. The second stage was observed at 182‐255 C and is due to the loss of two coordinated methanol molecules and one (CH3S)2 molecule (weight loss; Calcd./Found%; 16.79/17.50%). Emara et al. / European Journal of Chemistry 6 (2) (2015) 107‐116 113 Table 6. Temperatures of decomposition and the kinetic parameters of complexes 2, 3, 5 and 7. Compound Step T (K) A (1/s) ΔE (kJ/mol) ΔH (kJ/mol) ΔS (kJ/mol.K) ΔG (kJ/mol) [Ni(HL)]2.2H2O First 350 3.122 16.88 13.96 ‐245 99.71 Second 576 1.8×105 76.90 72.11 ‐158 163.29 Third 674 2291 65.91 60.30 ‐196 192.42 Fourth 937 9.577 53.61 45.81 ‐244 274.43 [Co(H2L)(H2O)2]2(NO3)2.2H2O First 364 1.8 19.11 16.08 ‐250 107.08 Second 562 1177 27.83 23.15 ‐200 140.23 Third 811 30.96 65.35 58.60 ‐233 247.86 [Fe(H2L)(CH3OH)NO3]2.2H2O.CH3OH First 441 0.07 10.91 7.24 ‐270 126.31 Second 516 23.82 33.13 28.83 ‐230 148.42 Third 655 3.90 20.37 14.92 ‐248 177.36 [Cd(HL)]2.H2O First 325 0.367 20.15 17.44 ‐260 101.94 Second 595 14.10 33.58 28.63 ‐245 174.41 Third 668 112 51.84 46.28 ‐220 193.24 The third stage was observed at 255‐800 C and is due to the dissociation of the organic ligand within the complex (weight loss; Calcd./Found%; 61.66/61.72%). The residue was iron oxide (weight loss; Calcd./Found%; 17.01/17.07%). Figure 2. TGA‐DrTGA of curves (A) [Fe(H2L)(CH3OH)NO3]2.H2O.CH3OH complex (5) and (B) [Cd(HL)]2.H2O complex (7). TGA of [Cd(HL)]2.H2O complex (7), shows three stages. Figure 2B depicts the TGA‐DrTGA curves of [Cd(HL)]2.H2O complex (7). The first stage is due to the loss of one uncoordinated water molecule (weight loss; Calcd./Found %; 2.15/1.76%) which observed at 50‐80 C. The second stage was observed at 247‐354 C and is due to the loss of two CH3SCN molecules and two CH3CN molecules (weight loss; Calcd./Found%; 27.31/26.08%). The third stage was observed at 354‐600 C and is due to loss of two CH3CHO molecules, SO2 and C6H4 molecules from the complex molecules (weight loss; Calcd./Found%; 27.31/27.63%). The residue was observed due to the decomposition of the rest part of the complex (weight loss; Calcd./ Found%; 43.60/44.00%). In order to access the influence of the type of the metal on the thermal behavior of the complexes, the order n, and the activation parameters of the various decomposition stages were determined from the TG thermograms using the Coats‐ Redfern equations in the following forms: ln[1‐(1‐α)1‐n /(1‐n)T2] = M/T +B for n ≠ 1 (1) ln[‐ln (1‐α) / T2] = M/T +B for n = 1 (2) where M = ‐E/R and B = ln AR/Ф E; E, R, A and Ф are the heat of activation, the universal gas constant, pre‐exponential factor and heating rate, respectively [55,56]. The correlation coefficient, r, was computed using the least square method for different values of n by plotting the left– hand side of Equation (1) or (2) versus 1000/T, Figure 3. The n values, which gave the best fit (r ≈ 1). From the intercept and linear slope of such stage, the A and E values were determined. The other kinetic parameters, ΔH, ΔS and ΔG were computed using the relationships; ΔH = E – RT, ΔS = R[ln(Ah/kT)‐1] and ΔG = ΔH ‐ TΔS, where k is the Boltzmann’s constant and h is the Plank’s constant. The kinetic parameters are listed in Table 6. The following remarks can be pointed out: (1) All complexes decomposition stages show a best fit for (n = 1) indicating a first order decomposition in all cases. (2) The value of ΔH increases significantly for the subsequently decomposition stages. (3) The negative values of activation entropies ΔS indicate a more order activated complex than reactants and/or the reaction are slow. (4) The positive values of ΔH means that the decomposition processes are endothermic [55,56]. Finally, on the basis of elemental analysis and infrared, mass, electronic spectra, ESR for copper(II) and oxovana‐ dium(IV) complexes and TGA for complex 2, 3, 5 and 7, it is possible to draw up the tentative structures of the transition metal complexes. Figure 4 depicts the proposed structures of the metal complexes. 3.3. Biological activity The hydrazone, H3L, ligand and its metal complexes were evaluated for sensitive organisms, which are: two strain Gram‐ positive bacteria (Staphylococcus aureus and Streptococcus pyogenes), Gram‐negative bacteria (Pseudomonas phaseolicola and Pseudomonas fluorescens) and two fungi (Fusarium oxysporum and Aspergillus fumigatus). The activity of the hydrazone, H3L, ligand and its metal complexes against the used organisms as depicted in Table 7. The hydrazone, H3L, ligand was found to be biologically active and some metal complexes exhibit mild antimicrobial activity. Although Ni(II) (2), Co(II) (3) complexes showed high antimicrobial activities, the Fe(III) complex (5) showed a dramatic low activity for all Gram‐positive, Gram‐negative bacteria and fungus strain. Cd(II) complex (7) showed high activity in Gram‐positive bacteria in high and low concentrations and low activity for Gram‐negative and fungi. Most of the complexes showed low activity towards the fungi. 114 Emara et al. / European Journal of Chemistry 6 (2) (2015) 107‐116 First stage of A First stage of B Second stage of A Second stage of B Third stage of A Third stage of B Figure 3. Costs‐Redfern plots for (A) [Fe(H2L)(CH3OH)NO3]2.H2O.CH3OH (5) and (B) [Cd(HL)]2.H2O (7). Figure 4. Structures of the metal complexes of H3L ligand. Emara et al. / European Journal of Chemistry 6 (2) (2015) 107‐116 115 Table 7. Biological activity of the synthesized hydrazone complexes 1‐7 a. Organisms Mean of zone diameter, nearest whole mm Gram‐positive bacteria Gram‐negative bacteria Fungi Staphylococcus aureus (ATCC 25923) Streptococcus pyogenes (ATCC 19615) Pseudomonas Phaseolicola (GSPB 2828) Pseudomonas fluorescens (S 97) Fusarium oxysporum Aspergillus fumigatus A B A B A B A B A B A B H3L 18‐I 14‐I 29‐H 22‐H 15‐I 15‐I 24‐I 18‐I 22‐I 12‐I 17‐I 13‐I [Cu(HL)]2 25‐I 16‐I 24‐I 14‐I 18‐I 13‐I 27‐H 16‐I 12‐L 13‐I 15‐I 9‐L [Ni(HL)]2.2H2O 30‐H 23‐H 26‐H 28‐H 32‐H 18‐H 28‐H 23‐H 14‐I 17‐I 17‐I 12‐I [Co(H2L)(H2O)2]2(NO3)2.2H2O 25‐I 23‐H 27‐H 26‐H 17‐I 22‐H 28‐H 17‐I 20‐I 14‐I 12‐L 8‐L [VO(HL)]2 22‐I 14‐I 15‐I 14‐I 16‐I 10I 13‐I 12‐I 15‐I 12‐I 8‐L 5‐L [Fe(H2L)(CH3OH)NO3]2.2H2O. CH3OH 19‐I 16‐I 13‐I 3‐L 4‐L 9‐L 6‐L 8‐L 8‐L 6‐L 2‐L 3‐L [Zn(HL)]2.6H2O 21‐I 15‐I 9‐L 14‐I 12‐I 14‐I 19‐I 15‐I 20‐I 10‐I 10‐I 9‐L [Cd(HL)]2.H2O 28‐H 22‐H 28‐H 24‐H 15‐I 22‐I 22‐I 14‐I 16‐I 12‐I 19‐I 15‐I Control b 42 28 38 30 36 25 38 30 40 28 40 31 a L: Low activity = Mean of zone diameter ≤ 1/3 of mean zone diameter of control. I: Intermediate activity = Mean of zone diameter ≤ 2/3 of mean zone diameter of control. H: High activity = Mean of zone diameter > 2/3 of mean zone diameter of control. b Cephalothin in the case of Gram‐positive bacteria, chloramphencol in the case of Gram‐negative bacteria and cycloheximide in the case of fungi. The concentration A is 2 mg/mL and B is 1 mg/mL. It is known that the chelation tends to make the ligand acts as more powerful and potent bacterial agent. A possible explanation for this increasing the activity upon chelation is that in chelating complex, the positive charge of the metal is partially shared with donor atoms present on the ligands and there is an electron delocalization over the whole chelating ring. This, in turn, increases the lipoid layers of the bacterial membranes. Generally, it is suggested that the chelating complexes deactivate various cellular enzymes, which play a vital role in various metabolic pathways of Other factors such as solubility, conductivity and dipole moment, which affected by the presence of metal ions, which may also be possible reasons for increasing the biological activity of the metal complexes as compared to the ligand from which they are derived [57,58]. 4. Conclusion Copper(II), nickel(II), cobalt(II), zinc(II), cadmium(II), iron(III) and oxovanadium(IV) complexes of hydrazone, H3L, ligand derived from the condensation S‐methyldithiocarbazate and 4,6‐diacetylresorcinol, in the molar ratio 1:1, has been synthesized. All the metal complexes were found to be dimmers. The structures of the ligand and its transition metal complexes were characterized by elemental analyses, spectral (infrared, electronic, mass, 1H NMR and ESR) data and magnetic susceptibility, molar conductivity measurements and thermal gravimetric analysis (TGA). The structure of the ligand is dibasic tridentate with ONS sites. The metal complexes exhibit different geometrical arrangements such as square planar, tetrahedral, square pyramidal and octahedral arrangements. 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