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Chapter 5 Complexes of Nickel(II), 3d8

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5.1. INTRODUCTION Nickel is one of only four elements that are magnetic at or near room temperature, the others being iron, cobalt and gadolinium. Its Curie temperature is 355°C, meaning that bulk nickel is non-magnetic above this temperature [1]. Some nickel compounds (nickel carbonyl, nickel sulphide) are considered to be highly toxic or carcinogenic. Pure nickel shows a significant chemical activity, though larger pieces of the metal are slow to react with air at ambient conditions due to the formation of a protective oxide surface. However, nickel is reactive with oxygen to the extent that native nickel is rare on Earth's surface, and is mostly confined to the interiors of larger nickel iron meteorites, which were protected from oxidation in space. Such native nickel is always found on Earth alloyed with iron, in keeping with the element's origin as a major end-product of the nucleosynthesis process, along with iron, in supernovas. An iron-nickel alloy is thought to compose the Earth's core. Due to slow rate of oxidation of nickel at room temperature, it is considered corrosion-resistant. Historically this has led to its use for plating metals such as iron and brass, to its use for chemical apparatus, and its use in certain alloys that will retain a high silvery polish, such as German silver. About 6% of world nickel production is still used for corrosion-resistant pure-nickel plating. Nickel was once a common component of coins, but has largely been replaced by cheaper iron for this purpose, especially since the metal has proven to be a skin allergen for some people. Although not recognized until the 1970s, nickel plays important roles in the biology of microorganisms and plants [2]. In fact, urease (an enzyme that assists in the hydrolysis of urea) contains nickel. The NiFe-hydrogenases contain nickel in addition to iron-sulphur clusters. Such [NiFe]-hydrogenases characteristically oxidise H2. A nickel-tetrapyrrole coenzyme, cofactor F430, is present in the methyl coenzyme Mreductase, which powers methanogenicarchaea. One of the carbon monoxide dehydrogenase enzymes consists of an Fe-Ni-S cluster [3]. Other nickel-containing enzymes include a class of superoxide dismutase [4] and a glyoxalase [5]. Uses for nickel complexes are virtually endless, but some discoveries of the 21st century exist that might be informative:

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• Scientist Daniel DuBois, from the Pacific Northwest National Laboratory, along with other scientists, discovered in 2011 that a nickel catalyst could help produce hydrogen for energy, quickly and efficiently. Up to this discovery, hydrogen came from platinum electrodes, an expensive and scare metal. By creating a nickel complex surrounded by pendant amines, proton relays channel proton into place so they can join to form hydrogen. This discovery may benefit the world as an abundant and inexpensive form of energy.

• Since renewable energy sources such as solar and wind power are intermittent, a way to store the energy is essential. Hydrogen gas, produced from nickel, could be the answer. The generation of hydrogen gas is efficient enough to convert all of the electricity captured through renewable sources into chemical energy. Replacing platinum catalysts with nickel complexes is a cost-effective solution. According to research teams from Pacific Northwest National Laboratory’s Center for Molecular Electrocatalysis and Villanova University continued to work on developing nickel-based complexes that increase the turnover frequency and decrease over-potential.

• Nickel complexes derived from nickel salt help adhere vulcanized rubber onto metal surfaces. Nickel salts suitable for adhesion are nickel and cobalt complexes along with sulfates, chlorides, bromides and acetates. A solution containing solvents, such as ethers, alcohols, acid amides, sulphones, sulphoxides, ketones or halogenated hydrocarbons dissolve the salts. When heated to above 212°C, the nickel mixture, along with a vulcanisable elastomer mixture, reacts. The reaction results in a mixture capable of permanently bonding to metal.

• Nickel organic complexes in agricultural film protect it from degradation due to UV light. Nickel ions stop the energy before it can break molecular bonds and generate free radicals. Some nickel complexes that used in polyolefin fiber applications, is resistant to pesticides, a desirable attribute in agriculture film. Polyolefin fiber is a light stabilizer used in plastics. Stabilizers containing nickel complexes pose potential toxicity concerns in the U.S., but the use of polyolefins continues due to their cost-effectiveness and performance.

Metal complexes of Ni(II) are also found to have DNA binding and photocleavage characteristics [11], act as a sensor [15], showed antibacterial activity against several

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pathogenic bacteria, such as Pseudomonas aeruginosa, Proteus vulgaris, Proteus mirabilis, Klebsiella pneumonia and Staphylococcus aureus [20], Escherichia coli, Staphylococcus aureus, Pseudomonas aeruginosa and Bacillus subtilis [22], Staphylococcus aureus, Hay bacillus, and Eschericha coli [23] and antifungal activities against several pathogenic fungi, such as Aspergillusniger and Pencillium chrysogenum [22]. In view of the above applications of Ni(II), the synthesis and characterization of the Ni(II) complexes with semicarbazide and thiosemicarbazide based ligands are highly desirable.

5.2. REVIEW OF LITERATURE The complexes of Fe(II) and Ni(II) with Schiff base derived from benzoin and 2-aminobenzoicacid have been prepared by Mohammed et al. [6]. Solubility, melting point, decomposition temperature, conductance measurement, infrared (IR) and UV-visible spectrophotometric studies were used in characterizing the compounds. The melting point of the Schiff base determined was 120oC. The decomposition temperatures of Fe(II) and Ni(II) complexes were 152oC and 155oC, while the molar conductance values were 11.2 and 10.7 ohm-1cm2mol-1, respectively. The UV-visible spectrophotometric analysis revealed 1:1 (metal-ligand) stoichiometry for the two complexes. DNA binding and photocleavage characteristics of a series of mixed ligand complexes of the type [M(phen)2LL]n+ (where M=Co(III), Ni(II) or Ru(II), LL=1,10-phenanthroline (phen), phenanthroline-dione (phen-dione) or dipyridophenazine (dppz) and n=3 or 2) have been investigated in detail by Arounaguiri et al. [7]. Various physico-chemical and biochemical techniques including UV/Visible, fluorescence and viscometric titration, thermal denaturation and differential pulse voltammetry have been employed to probe the details of DNA binding by these complexes; intrinsic binding constants (Kb) have been estimated under a similar set of experimental conditions. Analysis of the results suggested that intercalative ability of the coordinated ligands varies as dppz>phen>phen-dione in this series of complexes while the Co(II) and Ru(II) complexes investigated in this study effect photocleavage of the supercoiled pBR 322

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DNA, the corresponding Ni(II) complexes were found to be inactive under similar experimental conditions. Results of detailed investigations carried out inquiring into the mechanistic aspects of DNA photocleavage by [Co(phen)2(dppz)]3+ have also been reported.

Figure 5.1. Structures of the [M(phen)2LL]n+ type complexes investigated inthe study. Yakhvarov et al. [8] The reaction of [NiBr2(bpy)2] (bpy=2,20-bipyridine) with organicphosphinic acids ArP(O)(OH)H [Ar=Ph, 2,4,6-trimethylphenyl (Mes), 9-anthryl(Ant)] leads to the formation of binuclear Ni(II) complexes with bridging ArP(H)O2- ligands (Figure 5.2). Crystal structures of the binuclear complexes [Ni2(µ-O2P-(H)Ar)2(bpy)4]Br2 where (Ar=Ph, Mes, Ant) have been determined. In each structure, the metal ions have distorted octahedral coordination and were doubly bridged by two arylphosphinato ligands. Magnetic susceptibility measurements have shown that these complexes display strong antiferromagnetic coupling between the two nickel atoms at low temperatures, apparently similar to binuclear Ni(II) complexes with bridging carboxylato ligands.

Figure 5.2.

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Chiral Ni(II) complexes of Schiff base of (S)-2-N-(N-benzylprolyl)aminobenzophenone and 5-amino-4,5-dihydro-3-(2,4,6-trimethylphenyl)isoxazol-5-carboxylic acid (6) was synthesized by Mičúch et al. [9] via the cycloaddition of chiral complex of Ni with mesitonitrile oxide. The cycloaddition proceeded with complete regioselectivity to provide 5-substituted isoxazolines. The approach of the dipole took place predominantly from the less sterically hindered side of dipolarophile, the diastereoisomers were formed in 96:4 ratio. The detailed structure of complex was established by X-ray analysis.

Figure 5.3. Crystelline structure of compound 6 with crystallographic numbering

Qin et al. [10] prepared Ni(II) and Zn(II) complexes with the two new chromenone-based Schiff-base ligands, 3-{[(1,5-dihydro-3-methyl-5-thioxo-4H-1,2,4-triazol-4-yl)imino]methyl}-6-hydroxy-4H-1benzopyran-4-one (L1) and 2,2’-bis[(6-hydroxy-4-oxo-4H-1-benzopyran-3-yl)methylene]carbonothioic dihydrazide (L2). All the complexes were characterized by elemental analysis, IR data, and molar conductivity. The binding of these four complexes to calf-thymus DNA was carefully investigated by UV-vis spectroscopy, fluorescence spectroscopy, viscosity measurements and CD spectra. The experimental results indicated that the four complexes bind to calf-thymus DNA in an intercalative mode, with theintrinsic binding constants (K) of 3.94×104 ([NiL1]), 5.15×103 ([ZnL1]), 4.12×104 ([NiL2]), and 3.75×104m-1 ([ZnL2]). These data show that the complexes of L2 can interact more strongly with DNA than complexes of L1, and the Ni(II) complexes have a higher binding constant than Zn(II) complexes.

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The two nickel chelates of Schiff bases, 3-hydroxy-N-{2-[(3-hydroxy-N-phenylbutyrimidoyl)-amino]-phenyl}-N'-phenylbutyramidine (M(1)) and bis-4-(ethyliminomethyl)naphthalene-1-ol (M(2)), have been synthesized and explored as ionophores for preparing PVC-based membrane sensors selective to nickel ion by Gupta et al. [11]. The influences of membrane compositions on the potentiometric response of the electrodes have been found to substantially improve the performance characteristics. The best performance was obtained with the electrode having a membrane composition (w/w; mg) of (M(1)):PVC:NaTPB:CN in the ratio 5:150:5:150. The sensor showed a linear potential response for Ni(II) over a wide concentration range 1.6x10-7 to 1.0x10-2 M with Nernstian compliance (30.0+/-0.2 mV/decade of activity) within pH range 2.5-9.5 and a fast response time of 10 s. The sensor has been found to work satisfactorily in partially non-aqueous media up to 20% (v/v) content of methanol, ethanol, and acetonitrile and could be used for a period of 4 months. The analytical usefulness of the proposed electrode has been evaluated by its application in the determination of nickel in real samples. The practical utility of the membrane electrode has also been observed in the presence of surfactants. Four Ni(II) complexes, [Ni2L2(NO2)2]·CH2Cl2·C2H5OH, 2H2O (1),[Ni2L2(DMF)2(µ-NO2)]ClO4·DMF (2a), [Ni2L2(DMF)2(µ -NO2)]ClO4 (2b) and[Ni3L’2(µ3-NO2)2(CH2Cl2)]n·1.5H2O (3) where HL=2-[(3-amino-propylimino)-methyl]-phenol, H2L’=2-({3-[(2-hydroxy-benzylidene)-amino]-propylimino}-methyl)-phenol and DMF=N,N-dimethylformamide, have been synthesized by Naiya et al. [12] starting with the precursor complex [NiL2]·2H2O, Ni(II) perchlorate and sodium nitrite and characterized structurally and magnetically. The structural analyses revealed that in all the complexes, Ni(II) ions possess a distorted octahedral geometry. Complex 1 was a dinuclear di-µ2-phenoxo bridged species in which nitrite ion acts as chelating co-ligand. Complexes 2aand 2b also consist of dinuclear entities, but in these two compounds a cis-(m-nitrito-1kO:2kN) bridgeis present in addition to the di- µ2-phenoxo bridge. The molecular structures of 2a and 2b were equivalent; they differ only in that 2a contains an additional solvated DMF molecule. Complex 3 was formed by ligand rearrangement and is a one-dimensional polymer in which double phenoxo as well asµ-nitrito-1kO:2kN bridged trinuclear units were linked through a very rare m3-nitrito-

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1kO:2kN:3kO’bridge. Analysis of variable-temperature magnetic susceptibility data indicated that there is a global weak antiferromagnetic interaction between the Ni(II) ions in four complexes, with exchange parameters J of -5.26, -11.45, -10.66 and -5.99 cm-1 for 1, 2a, 2b and 3, respectively.

Figure 5.4. Structure of the Ni(II) environments in compounds 1–2(2a and 2b) with the bridging bond angle (°) and coupling constant (cm-1). Color code: N =blue, O = red, Ni = green.

Complexes of Cu(II), Ni(II), Mn(II), Zn(II) and VO(II) with neutral tetradentate N2O2

have been synthesized by Raman et al. [13] using a Schiff base formed by the condensation of o-phenylenediamine with acetoacetanilide in alcohol medium. All complexes were characterized on the basis of their microanalytical data, molar conductance, magnetic susceptibility, IR, UV–Vis 1H NMR and ESR spectra. IR and UV–vis spectral data suggested that all the complexes are square-planar except the Mn(II) and VO(II) chelates, which are of octahedral and square pyramidal geometry respectively. The monomeric and neutral nature of the complexes was confirmed by their magnetic susceptibility data and low conductance values. The ESR spectra of copper and vanadyl complexes in DMSO solution at 300K and 77 K were recorded and their salient features were reported.

Ercag et al. [14] have prepared Schiff base ligand (HL) derived from 4-hexylaniline with isatin (1H-indole-2,3-dione)and its complexes with Cu(II), Ni(II) and characterized by analytical, spectroscopic (IR, UV-vis, Mass) techniques, electrical conductivity, magnetic and thermal measurements. The crystal and molecular structure of [Cu(HL)2Cl2] was determined by a single-crystal X-ray diffraction study. The molecular structure of the title compound has an inversion center on the Cu atom.

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A series of Werner complexes featuring the tridentate ligand smif, that is, 1,3-di-(2-pyridyl)-2-azaallyl, have been prepared by Frazier et al. [15]. Synthesis of (smif)2M (1-M; M=Cr, Fe) were accomplished via treatment of M(NSiMe3)2(THF)n (M=Cr, n=2; Fe, n=1) with 2 equiv of (smif)H (1,3-di-(2-pyridyl)-2-azapropene); ortho-methylated (Mesmif)2Fe (2-Fe) and(Me2smif)2Fe (3-Fe) were similarly prepared. Metatheses of MX2 variants with 2 equiv of Li(smif) or Na(smif) generated 1-M (M = Cr, Mn, Fe, Co, Ni, Zn, Ru). Metathesis of VCl3(THF)3 with 2 Li(smif) with a reducing equiv of Na/Hg present afforded 1-V, while 2 Na(smif) and IrCl3(THF)3 in the presence of NaBPh4 gave [(smif)2Ir]BPh4 (1+-Ir). Electrochemical experiments led to the oxidation of 1-M (M=Cr, Mn, Co) by AgOTf to produce [(smif)2M]OTf (1+-M), and treatment of Rh2(O2CCF3)4 with 4 equiv Na(smif) and 2 AgOTf gave 1+-Rh. Characterizations by NMR, EPR, and UV−vis spectroscopies, X-ray crystallography and DFT calculations were presented. Intraligand (IL) transitions derived from promotion of electrons from the unique CNCnb (nonbonding) orbitals of the smif backbone to ligand π*-type orbitals are intense (ε≈10 000−60 000 M−1cm−1), dominate the UV−visible spectra, and give crystals a metallic-looking appearance. The Cu(II), Ni(II), Fe(II) and Zn(II) Schiff base complexes were prepared from salicylaldehyde and o-amino benzoic acid by Johari et al.[16]. The chemical structure of the synthesized metal ligand complexes were confirmed IR and NMR spectral analysis. The free Schiff base and its complexes have been tested for their antibacterial activity against several pathogenic bacteria, such as Pseudomonas aeruginosa, Proteus vulgaris, Proteus mirabilis, Klebsiella pneumonia and Staphylococcus aureus. The antibacterial activity was determined by the Agar Ditch technique using DMF (polar) and 1, 4 dioxane (non polar) as solvent. The metal complexes showed differential effect on the bacterial strain. The antibacterial activity is dependent on the molecular structure of the compound, the solvent used and the bacterial strain under consideration. In the non polar solvent 1,4-dioxane the best antibacterial activity was shown by the Zn complex while in polar solvent DMF, Ni complex of Schiff base showed best antibacterial activity. A bis(hydrazone) was prepared by reacting isatinmonohydrazone with 2-hydroxy-1-naphthaldehyde and a series of metal complexes with this new ligand were synthesized

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by reaction with Mn(II), Fe(II), Co(II), Ni(II), Cu(II) and Zn(II) salts by Murukan et al. [17]. These complexes were characterized on the basis of elemental analyses, molar conductance, magnetic susceptibility data, UV-vis, IR, ESR, and NMR spectral studies, wherever possible and applicable. Analytical data reveal that the Ni(II), Cu(II) and Zn(II) complexes possess 1:1 metal–ligand ratios and that Mn(II), Fe(II) and Co(II) complexes exhibit 1:2 ratios. Infrared spectral data suggested that the bis(hydrazone) behaves as a monobasic tridentate ligand with ONO donor sequence towards the metal ions. X-ray diffraction study of the Cu(II) complex indicated an orthorhombic crystal lattice. The EPR spectral data showed that the metal–ligand bond has considerable covalent character. The electrochemical behaviour of the Cu(II) complex was investigated by cyclic voltammetry (CV). Antibacterial tests of the ligand and the metal complexes were also carried out and it has been observed that the complexes are more potent bactericides than the ligand. A series of first complexes of Co(II), Ni(II), Cu(II), Mn(II) and Fe(III) have been synthesized with Schiff base derived from isatinmonohydrazone and fluvastatin by Kulkarni et al.[18]. These complexes have been characterized in the light of elemental analyses, spectral (IR, Uv-vis., FAB-mass and ESR) and magnetic studies. The elemental analyses of the complexes confine to the stoichiometry of the type ML2.2Cl [M=Co(II), Ni(II), Cu(II) and Mn(II)] and [FeL2.2Cl]Cl. The redox properties of the complexes were extensively investigated by electrochemical method using cyclic voltammetry (CV). The Co(II) and Cu(II) complexes exhibited quasi-reversible single electron transfer process whereas Mn(II) and Fe(III) complexes shows two redox peaks of quasi-reversible one electron transfer process. The Schiff bases and their complexes have been screened for their in-vitro antibacterial (Escherichia coli, Staphylococcus aureus, Pseudomonas aeruginosa and Bacillus subtilis) and antifungal (Aspergillus niger and Pencillium chrysogenum) activities by minimum inhibitory concentration (MIC) method. A thiosemicarbazone derivative of the Vitamin K3 has been synthesized by Tang et al. [19]. Five transition metal complexes (Mn(II), Co(II), Ni(II), Cu(II) and Zn(II) complexes) of the thiosemicarbazone have been prepared and characterized by IR, UV-vis, molar conductance and thermal analyses. All of these complexes possess strong

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inhibitory action against Staphylococcus aureus, Hay bacillus, and Eschericha coli. The antibacterial activities of the complexes were stronger than those of the ligand. Co(II), Ni(II) and Cu(II) complexes were synthesized with thiosemicarbazone (L1) and semicarbazone (L2) derived from 2-acetyl furan by Chandra et al. [20] These complexes were characterized by elemental analyses, molar conductance, magnetic moment, mass, IR, electronic and EPR spectral studies. The molar conductance measurement of the complexes in DMSO corresponds to non-electrolytic nature. All the complexes are of high-spin type. On the basis of different spectral studies six coordinated geometry may be assigned for all the complexes except Co(L)2(SO4) and Cu(L)2(SO4) [where L=L1 and L2] which were of five coordinated square pyramidal geometry. Beraldo et al. [21] prepared the complexes of Co(II) and Ni(II) with N(4’)-methyl and N(4’)-ethyl thiosemicarbazones derived from 3- and 4-acetylpyridine and characterized by microanalyses, magnetic susceptibility and molar conductivity measurements and by their electronic, IR and NMR (in the case of Ni(II)complexes) spectra. Asmy et al. [22] have investigated the ligation behavior of 2-hydroxybenzophenone and 2-hydroxy-4-methoxybenzophenone N substituted thiosemicarbazones towards Ni(II) and Cu(II) ions. The isolated complexes were identified by elemental analyses, molar conductance, magnetic moment, IR, UV–vis and ESR spectral studies. The IR spectra indicated that the investigated thiosemicarbazones lost the N2 proton or the N2 and OH protons and act as mononegative or binegative tridentate ligands. The ligands containing methoxy group facilitate the deprotonation of OH by resonance more than the SH. Most of the Ni(II) complexes measured subnormal magnetic moments due to square-planar + tetrahedral conFigureuration and supported by the electronic spectra. The percentage of square-planar to tetrahedral was calculated and found in agreement with the ligand splitting energy (10Dq). Chandra et al. [23] synthesized the Ni(II) complexes having the general composition Ni(L)2X2 [where L:isopropyl methyl ketone semicarbazone (LLA), isopropyl methyl ketone thiosemicarbazone (LLB), 4-aminoacetophenone semicarbazone (LLC) and 4-aminoacetophenone thiosemicarbazone (LLD) and X=Cl−, ½SO42−]. All the Ni(II) complexes reported have been characterized by elemental analyses, magnetic moments,

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IR, electronic and mass spectral studies. All the complexes were found to show magnetic moment corresponding to two unpaired electrons. The possible geometries of the complexes were assigned on the basis of electronic and infrared spectral studies. Newly synthesized ligand and its Ni(II) complexes have been screened against different bacterial and fungal growth.

5.3. PRESENT WORK In the present chapter we are discussing the synthesis and characterization of Ni(II) complexes with six semicarbazide and thiosemicarbazide based ligands (L1, L2, L3, L4, L5 and L6). These complexes were characterized by elemental analyses, molar conductance, magnetic moment, electronic and IR spectral studies. PREPARATION OF NICKEL(II) COMPLEXES The complexes with ligands L1, L2, L3, L4, L5 and L6 were prepared by using semicarbazide and thiosemicarbazide based ligands and corresponding Ni(II) salts [NiX2.nH2O] (Where X = Cl-, NO3-,½ SO42- and CH3COO-). Synthesis of complexes with ligands L1- L6: A hot ethanolic (or aqueous in case of sulphate salt) (20 mL) solution of corresponding metal salts (0.001 mol) was mixed with hot ethanolic solution of the respective ligands (L1- L6) (0.002 mol). The mixture was refluxed at 80°C (±5). The conditions for the different metal complexes are given in the table 5.1. On cooling the contents, colored complexes were precipitate out. These were filtered, washed with 50% ethanol and dried in vacuum over P4O10.

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Table 5.1. Reaction conditions for different metal complexes

Complex Reflux time (h) pH Color

Melting Point / Decomposition Temp.

(°C) [Ni(L1)2Cl2] 10 8 Green 240 [Ni(L1)2(NO3)2] 14 8 Green 250 [Ni(L1)2SO4] 15 6.5 Brown 210 [Ni(L1)2(CH3COO)2] 12 8 Dark green 240 [Ni(L2)2Cl2] 4 6.5 Green Above 270 [Ni(L2)2(NO3)2] 40 6.5 Green Above 270 [Ni(L2)2SO4] 14 6.5 Brown Above 270 [Ni(L2)2(CH3COO)2] 1 6.5 Brown Above 270 [Ni(L3)2Cl2] 6 6 Green Above 270 [Ni(L3)2(NO3)2] 18 8 Green Above 270 [Ni(L3)2SO4] 14 8 Green Above 270 [Ni(L3)2(CH3COO)2] 13 6.5 Green Above 270 [Ni(L4)2Cl2] 23 6 Green Above 270 [Ni(L4)2(NO3)2] 22 8 Dark Green Above 270 [Ni(L4)2SO4] 36 5 Green Above 270 [Ni(L4)2(CH3COO)2] 13 5 Green Above 270 [Ni(L5)2Cl2] 21 7.5 Green Above 270 [Ni(L5)2SO4] 23 8 Brown Above 270 [Ni(L6)2Cl2] 8 7.5 Green Above 270 [Ni(L6)2SO4] 10 7.5 Dark green Above 270 [Ni(L6)2(CH3COO)2] 8 6 Brown Above 270 5.4. RESULTS AND DISCUSSION On the basis of elemental analysis the complexes were found to have general compositions [Ni(L)2X2] (where L=L1-L6, X=CI¯, NO3-, ½ SO42-, CH3COO-]. The molar conductance data of these complexes in dimethylsulphoxide (DMSO) indicated that these complexes are non-electrolytes in nature. These complexes may be formulated as [Ni(L)2X2][X= CI¯, NO3-, ½ SO42-, CH3COO-] (Table 5.2).

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Table 5.2. Molar conductance and elemental analyses of Ni(II) complexes

Complex M.wt. Molar conductance

Yield (%)

Elemental analysis Found (Calculated) M C H N

[Ni(L1)2Cl2] 639.7 11.1 59 9.06 33.45 3.02 13.05 NiC18H20N6O2Br2Cl2

(9.18) (33.77) (3.13) (13.13)

[Ni(L1)2(NO3)2] 692.7 10.4 57 8.12 31.01 2.58 16.04 NiC18H20N8O8Br2

(8.47) (31.18) (2.89) (16.17)

[Ni(L1)2SO4] 664.7 12.6 60 8.76 32.43 2.92 12.53 NiC18H20N6O6Br2S

(8.83) (32.50) (3.01) (12.64)

[Ni(L1)2(CH3COO)2] 686.7 11.6 61 8.26 38.19 3.57 12.04 NiC22H26N6O6Br2

(8.55) (38.44) (3.79) 12.23

[Ni(L2)2Cl2] 671.7 10.2 58 8.64 32.01 2.91 12.35 NiC18H20N6O2Br2S2Cl2

(8.71) (32.06) (2.97) (12.47)

[Ni(L2)2(NO3)2] 724.7 13.2 53 8.02 29.73 2.64 15.32 NiC18H20N8O6Br2S2

(8.10) (29.81) (2.76) (15.45)

[Ni(L2)2SO4] 696.7 10.4 56 8.32 31.09 2.75 12.07 NiC18H20N6O4Br2S3

(8.40) (30.91) (2.86) (12.02)

[Ni(L2)2(CH3COO)2] 718.7 13.4 54 8.02 36.52 3.48 11.37 NiC22H26N6O4Br2S2

(8.17) (36.73) (3.62) (11.69)

[Ni(L3)2Cl2] 535.9 12.1 57 10.79 49.11 4.81 15.52 NiC22H26N6O2Cl2

(10.96) (49.28) (4.85) (15.68)

[Ni(L3)2(NO3)2] 588.7 10.8 58 9.86 44.78 4.31 18.89 NiC22H26N8O8

(9.97) (44.84) (4.42) (19.02)

[Ni(L3)2SO4] 560.7 11.7 61 9.78 46.94 4.52 15.87 NiC22H26N6O6S

(10.47) (47.08) (4.64) (14.98)

[Ni(L3)2(CH3COO)2] 582.7 12.7 55 9.89 53.43 5.32 14.30 NiC26H32N6O6

(10.07) (53.54) (5.49) (14.42)

[Ni(L4)2Cl2] 567.7 12.1 56 10.21 46.43 4.42 14.73 NiC22H26N6S2Cl2

(10.34) (46.50) (4.58) (14.80)

[Ni(L4)2(NO3)2] 620.7 13.6 57 9.34 42.41 4.08 17.92 NiC22H26N8O6S2

(9.46) (42.53) (4.19) (18.04)

[Ni(L4)2SO4] 592.7 14.4 56 9.78 44.48 4.28 14.02 NiC22H26N6O4S3

(9.90) (44.54) (4.39) (14.17)

[Ni(L4)2(CH3COO)2] 614.7 15.2 58 9.43 50.66 5.13 13.58 NiC26H32N6O4S2

(9.55) (50.76) (5.21) (13.67)

[Ni(L5)2Cl2] 691.7 13.2 54 8.38 55.46 4.27 12.08 NiC32H30N6O4Cl2

(8.49) (55.52) (4.34) (12.14)

[Ni(L5)2SO4] 716.7 14.8 57 7.98 53.46 4.08 11.64 NiC32H30N6O8S

(8.19) (53.58) (4.19) (11.72)

[Ni(L6)2Cl2] 723.7 12.8 55 8.02 52.97 4.01 11.54 NiC32H30N6O2S2Cl2

(8.11) (53.06) (4.15) (11.61)

[Ni(L6)2SO4] 748.7 10.6 61 7.76 51.11 3.93 11.17 NiC32H30N6O6S3

(7.84) (51.29) (4.01) (11.22)

[Ni(L6)2(CH3COO)2] 770.7 11.6 59 7.57 55.94 4.58 10.90 NiC36H36N6O6S2 (7.62) (56.05) (4.67) (10.81)

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MAGNETIC MOMENT The magnetic moment of octahedral or tetrahedral complexes lies in the range of 2.8-3.5 B.M. corresponding to two unpaired electrons. It is close to spin only value. But the four coordinate, square-planar complexes are diamagnetic.The ground state of a regular octahedral complex is 3A2g and no singlet levels arising from d configuration can cross it. The major contribution to the magnetic susceptibility of octahedral Ni(II) complex is given by the spin only term

S(S+1)N β2g2/3kT Where N= Avagadro number, β =Bohr magneton, g=2.00 and, S=1 There is small contribution because of spin-orbit coupling between the first excited state3T2g and the ground state 3A2g. It leads to a magnetic moment value of µef\f=g(1-4λ/10Dq),i.e. about 10% above the spin only value, where λ is the spin orbit coupling constant (-315cm-1) for the free ion Ni(II) and 10 Dq is the energy separation between the interaction level. if α is the mixing coefficient of the dx2-y2 and dz2 orbital in the antibonding molecular orbital we get g=2(1-4ά2/10Dq). There is also orbital contribution by temperature independent paramagnetism (T.I.P) arising from the first excited state 3T2gwhen is equal to 8NB2/10Dq [24]. Thus magnetic susceptibility of octahedral Ni(II) complexes can be obtained from following relation.

H = 8Nβ2/3kT+ (1-4λ2d/10Dq)2+8Nβ2/10Dq In fact, the value of the magnetic moment usually found between 2.9-3.3 B.M. and is almost independent of temperature. Regular or nearly regular tetrahedral complexes show characteristic magnetic and spectral properties. In Td symmetry, the d8configuration gives rise to a 3T1(F) ground state. The ground state term 3T1(F) has much inherent orbital angular momentum. The magnetic moment of truly tetrahedral Ni(II) should be ~4.2 B.M. at room temperature. However, even slight distortion reduces this markedly (by splitting the orbital degeneracy). Actually in tetrahedral field the stark pattern is inverted and triply degenerate level is the lowest one, which is further split by small rhombic field into three levels separated by energy interval, kT. Each of these sublevels is occupied by an extent dependent upon the relative size of kT and of the energy intervals. The orbital contribution is no longer small and the moments are expected to be somewhere between

Nickel(II), 3d8

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the spin only value of 2.83 B.M. and spin only plus the orbital contribution value of 4.2 B.M. [25]. Experimentally it has been found that tetrahedral Ni(II) complexes have moments in the range 3.4-4.2 B.M. However, the distortion in the octahedral field enhances the value of the magnetic moments. Figgis and Nyholm [26] suggested that distinction based on this is not a conclusive proof for deciding between the two geometries because the orbital contribution also depends upon the electronegativities [27] of the coordinated ligands and is very sensitive to the slight distortion from the cubic symmetry of the octahedral complexes. The values are in accordance with a high-spin configuration and show the presence of a distorted octahedral environment around the metal ions [28]. All the electrons in the 3d orbital are paired in four coordinate square planar complexes thus these complexes are diamagnetic. Magnetic moment gives the idea of geometry of Ni(II) complexes under study. Magnetic moment of complexes of present study lie in the range of 2.95-3.03 B.M., (Table 5.4) which correspond to two unpaired electrons. It indicated that these complexes may have six coordinate octahedral geometry and the complexes are having molecular formula [NiLX2]. However, it is quite difficult to assign the geometry only on the basis of magnetic moment. The possible geometry will be elucidated on the basis of electronic spectral studies. FT-IR FT-IR spectra of the ligands have already been explained in chapter 2. On complexation the position of bands due to ν(>C=N) and ν(>C=O) in semicarbazone complexes and due to ν(>C=N) and ν(>C=S) in case of thiosemicarbazone complexes, is shifted by 10-90 cm-1. This indicated that the coordination takes place through the nitrogen atoms of imine group, oxygen atom of the >C=O group in case of semicarbazone complexes and nitrogen atoms of imine group and sulphur atom of the >C=S group in case of thiosemicarbazone complexes and all the ligands behaved as bidentate (Table 5.3) (Figures 5.5 -5.19).

Nickel(II), 3d8

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Table 5.3. Important infrared spectral bands (cm−1) and their assignments.

Compounds Assignments Remarks ν(N−H) ν(C=O) ν(C=N) ν(C=S) ν(M−N) Ligand(L1)

3197

1708

1583

Ni(L1)Cl2 3195 1686 1570 437 [Ni(L1)(NO3)2] 3194 1676 1567 454 ν5=1410, ν1=1320,

ν2=1021 and ∆(ν5- ν1) =90 cm-1 indicating monodentate nature of nitrate group. [29]

Ni(L1)SO4 3288 1630 1509 442 ν3 splitted at 1149, 1106, 1090 and1032 cm-1 while ν1 at 994 cm-1 indicating bidentate nature of SO4

2- group[31].

[Ni(L1)(CH3COO)2] 3286 1658 1545 482 νas(OAc)=1490, νs(OAc) =1340, ∆ν=150 cm-1 indicating monodentate nature of acetate group. [29].

Ligand(L2) 3142 1585 852 Ni(L2)Cl2 3253 1438 702 435 [Ni(L2)(NO3)2] 3245 1457 770 487 ν5=1390, ν1=1280,

ν2=1040 and ∆(ν5 - ν1)=110 cm-1 indicating monodentate nature of nitrate group.

Ni(L2)SO4 3120 1654 711 455 ν3 splitted at 1100, 1063, and 1035 while ν1 at 971 cm-1 indicating bidentate nature of SO4

2- group. [Ni(L2)(CH3COO)2] 3134 1508 737 440 νas(OAc)=1475,

νs(OAc)=1307, ∆ν=168 cm-1 indicating monodentate nature of acetate group.

Ligand(L3) 3180 1690 1586 Ni(L3)Cl2 3178 1643 1550 451 [Ni(L3)(NO3)2] 3175 1686 1570 533 ν5=1450, ν1=1380,

ν2=1070 and ∆(ν5 - ν1)=70 cm-1 indicating monodentate nature of nitrate group.

Ni(L3)SO4 3184 1622 1578 434 ν3 splitted at 1153, 1095, and 1056 while ν1 at 976 cm-1 indicating bidentate nature of SO4

2- group.

Nickel(II), 3d8

160

Compounds Assignments Remarks ν(N−H) ν(C=O) ν(C=N) ν(C=S) ν(M−N) [Ni(L3)(CH3COO)2]

3176

1654

1567

459

νas(OAc)=1430, νs(OAc)=1257, ∆ν = 183 cm-1 indicating monodentate nature of acetate group.

Ligand(L4) 3142 1600 846 Ni(L4)Cl2 3130 1590 789 435 [Ni(L4)(NO3)2] 3150 1609 720 480 ν5 =1370, ν1=1280, ν2

=1099, ∆(ν5 - ν1)=90 cm-1 indicating monodentate nature of nitrate group

Ni(L4)SO4 3368 1547 766 438 ν3splitted at 1160, 1120, 1090 and 1054 while ν1 at 956 cm-1 indicating bidentate nature of SO4

2- group

[Ni(L4)(CH3COO)2] 3369 1547 763 426 νas(OAc)=1448, νs(OAc)=1291, ∆ν = 157 cm-1 indicating monodentate nature of acetate group.

Ligand(L5) 3147 1693 1571 Ni(L5)Cl2 3195 1643 1547 476 Ni(L5)SO4 3193 1687 1458 480 ν3splitted at 1149, 1115

and 1069 while ν1 at 990 cm-1 indicating bidentate nature of SO4

2- group

Ligand(L6) 3144 1600 762 Ni(L6)Cl2 3143 1518 757 449 Ni(L6)SO4 3145 1590 760 490 ν3 splitted at 1190, 1130

and 1083 while ν1 at 990 cm-1 indicating bidentate nature of SO4

2- group

[Ni(L6)(CH3COO)2] 3141 1520 757 440 νas(OAc)=1452, νs(OAc)=1313, ∆ν=140 cm-1 indicating monodentate nature of acetate group.

Nickel(II), 3d8

161

Figure 5.5. FT-IR spectra of [Ni(L1)2Cl2]

Figure 5.6. FT-IR spectra of [Ni(L1)2SO4]

Nickel(II), 3d8

162



Nickel(II), 3d8

163

Figure 5.9. FT-IR spectra of [Ni(L2)2(CH3COO)2]


Nickel(II), 3d8

164


Figure 5.12. FT-IR spectra of [Ni(L3)2(NO3)2]

Nickel(II), 3d8

165


Figure 5.14. FT-IR spectra of [Ni(L4)2(NO3)2]

Nickel(II), 3d8

166



Nickel(II), 3d8

167



Nickel(II), 3d8

168

Figure 5.19. FT-IR spectra of [Ni(L6)2(CH3COO)2]

ELECTRONIC SPECTRA The Ni(II) ion has 3d8 valance electron configuration. This gives rise to the Russel-Saunders terms (in order of increasing energy) 3F, 1D, 3P, 1G, 1S. The triplet term, 3F4 represents the electronic ground state terms of the gaseous ion. Ni(II) complexes are known with a six-coordinate octahedral, with five-coordinate square pyramidal or trigonal pyramidal and with four-coordinate square planar or tetrahedral configurations. It is a peculiarity of Ni(II) chemistry that complexes of one configuration can be easily converted to other configurations. This structural liability implies that the free energy difference between different stereo chemical forms is usually small. The Ni(II) complexes with coordination number six in almost all cases have high-spin electronic configuration and regular or distorted stereochemistry. The six coordinated complexes have either Oh or D4h symmetry. In octahedral field the strong field

Nickel(II), 3d8

169

configuration is t2g6 eg2 (S=1) follow the same pattern of splitting in electronic spectra as d3 with 3A2g lying as ground state and the correlation diagram is directly analogous to that of d3. Thus three spin allowed transitions are as follow:

3A2g(F)→3T2g(F) (v1)=10Dq 3A2g(F)→3T1g(F) (v2)=18Dq 3A2g (F)→3T1g (P) (v2)=12Dq + 15B

Due to ‘non-crossing’ rule v2 cannot lie higher in energy than v3. Since d8 has 3A2g ground state in both weak and strong fields, thus the term diagram is similar in both the cases. These bands occur in the range of 6500-13000, 11000-20000 and 19000-27000 cm-1 region respectively [30]. Considering crystal field theory the Ni(II)complex in case of Oh symmetry splits into triplet terms 3F and 3P. However, the energy level in octahedral field having different configurations can be represented as: t2g6eg23A2-12Dq 1E-12Dq + 8B + 2C – 6B2/10Dq 1A1-12Dq + 16B + 4C – 10B2/10Dq t2g5eg33T2-2Dq 3T17.5B + 3Dq – ½(225B2 + 100Dq2 – 180qB)1/2 1T2-2Dq + 8B + 2C – 12B2/10Dq 1T2-2Dq + 12B + 2C t2g4eg43T17.5B + 3Dq + ½(225B2 + 100Dq2 – 180qB)1/2

If the six coordinated Ni(II) complexes possesses D4h symmetry and the anions are at axial positions, the splitting will be as follows [31].

.

Nickel(II), 3d8

170

3T1g(P)

3Egc

3A2gb

3Ds-5Dt

3T1g(F)

3T2g(F)

3A2g(F)

3Egb

3A2ga

3B2g

3Ega

3B1g

-6Ds - 5/4Dt

35/4Dt

10Dqxy

If the complexes are four coordinated then it can have either tetrahedral or square planar geometry. Diamagnetic square-planar Ni(II) complexes having D4h symmetry, the five degenerated d-orbital are splitted as, dx2-y2 (b1g), dxy (b2g), dz2 (a1g), dxz, dyz (eg). The electronic ground state of square planar complex may be either (eg)4 (a1g)2 (b2g)2, a spin-singlet state, term 2A1g or (eg)4(a1g)2(b2g)1(b1g)1, a spin-triplet state, 3A2g and an excited state 1A2g. The relative stability of the 1A1g and 3A2g states is determined by the energy separation of the dxy and dx2-y2orbitals. The majority of square-planarNi(II) complexes exhibits three bands in the region of 15000-180000, 20000-25000 and near 30000 cm-1 assignable to the 1A1g(D) →1A2g(G), 1A1g(D)→1B2g(G), 1A1g(D)→1Eg(G) transitions respectively. Weak bands in the near infrared region have been assigned as spin forbidden transitions. For the five-coordinate complex, both low-spin and high-spin complex are known. They either possess C4v symmetry (square-pyramidal) and D3h symmetry (trigonal pyramidal). The free ion term 4F split into five levels in a field of C4v symmetry and four levels in a field of D3h symmetry. Thus the d-orbital split as:

Nickel(II), 3d8

171

(b)dx2-y2

(a1)dz2

(b2)dxy

(e)dxy, dyz

d

C4v D3h

dz2(a1)

dx2-y2, dxy(e')

dxz, dyz(e")

For square-pyramidal Ni(II) complexes six-spin allowed transitions are possible[31]. One of these 3B1 (F)→3A2 (P) is a two electrons transition in the strong field limit. The transitions to the states 3E (F) and 3E (P) can acquire intensity by the mixing of metal d and p orbital. The spectrum exhibits four bands in the range 6,000 -20,000 cm-1. The first and fourth bands are more intense than the others [32]. On the basis of above discussion it is easy to confirm the geometry of complexes under study. Electronic spectra of six-coordinate Ni(II) complexes may possess Oh or D4h symmetry. From the energy level diagram, for octahedral complexes the three spin allowed transitions should be observed. The electronic spectra of the complexes [NiLX2] (where X = Cl-, NO3-, SO42-, CH3COO-, show three bands in the range of 9294-10173 cm-1 and 10152-17094 cm-1 and 20576-30211cm-1 [Figures 5.20-5.32] assignable to 3A2g(F) →3T1g(F) (v1), 3A2g(F) →3T2g(F) (v2) and 3A2g(F) →3T1g(P) (v3) transitions respectively. There is no splitting of the v1 band, which is generally observed in D4h symmetry. Thus complexes under study possess distorted octahedral geometry. LIGAND FIELD PARAMETERS The ligand field parameters have been calculated and are listed in Table 5.4. The triplets terms 3F and 3P arising from 3d8configuration of the free ion Ni(II) (which have an energy separation equal to 15B in terms of the reach interelectronic repulsion parameters) are represented as:

Nickel(II), 3d8

172

3A2g(F) E = -12Dq (a) 2T2g(F) E = -2Dq (b) 2T1g(F) and 3T1g(P): [6DqP – 16(Dq)2] + [-6Dq-P]E =E2 = 0 (c)

where P is the energy of the 3P state. From equation (a) and (b) it is seen that the energies of both 3A2g(F) and 3T2g(F) are linear function of Dq. The difference in transition of 3T2g state and 3A2g state is directly corresponds to Dq. Considering the energy of other states equation (c) can be solved. However, this equation has been derived by assuming that the ligands are point charges or point dipoles and there is no covalence in the metal-ligand bond. If this was true, the value for Dq just determine could be substituted into equation (c), the energy of 3P obtained from the atomic spectrum of the gaseous ion and the energy of the two levels in the complex calculated from the equation (c). The frequency of the expected spectral transitions are calculated for one band corresponding to the difference between the energy of the levels 3A2g(F) �3T2g(F) and for the other band from the energy difference3A2g(F) �3T1g(P). The experimental energies obtained from the spectra almost always lower than the values calculated in this way. The deviation attributed to covalency in the bonding. The effect of the covalency is to reduce the positive charge on the metal ion, as a consequence of the inductive effect of the ligands. With reduced positive charge, the radial extension of the d orbitals increasing, this decreases the electron-electron repulsion and lowering the energy of the 3P states. Covalency is foreign to the crystal field approach and is incorporated into the ligands-field approach by providing additional parameters. The difference in energy between 3P and 3F states in the complex related to that in the gaseous ion is decreased by the covalency and as result, the gas phase value cannot be used for P in the equation (c); rather, P must be experimentally evaluated for each complex. Equation (c) can be calculated by using the Dq value from the transitions 3A2g(F) �3T2g(F) and the experimental energy, ∆E from the 3A2g(F) �3T1g(P) transition. Thus from P can be calculated. The lowering of 3P is a measure of covalency other among effects. It is referred to as the nephelauxetic effect and is sometimes expressed by a parameter �, a percentage lowering of the energy of 3P state in the complex

Nickel(II), 3d8

173

compared to the energy 3P in the free gaseous ion. It is calculated by using equation as follows:

β = [B-B’]/B x 100 (d) where B is the Racah parameter discussed earlier for the gaseous ion. It should be noted that P of equation (c) proportional to B. in the case of Ni(II), the energy of 3P in complex can be substituted along with Dq into equation (c) and the other root calculated. The difference in the energy between this root and the energy of 3A2g gives the frequency of middle band 3A2g→3T1g(F). The agreement of the calculated and experimental values for this band is good evidence for D4h symmetry. The value of B for a given complex can be calculated as:

B = (v2 + v3 – 3v1) / 15 (e) The value of Racah parameter B is found in complexes is less than the value of 1041 cm-1 found in the free ion. The value of Dq found in octahedral complexes of Ni(II) vary between 968-1089 cm-1, depending on the position of the ligand in the spectrochemical series. The β is defined as:

β = B (complex) / B(free ion) The value of β depends on the position of the ligand in the spectrochemical series. The value of β lies in the range of 0.18-0.98. The two quantities are easily related if equation (d) is rewritten as:

β = (1-B’) x 100

Nickel(II), 3d8

174

Figure 5.20. UV-visible spectrum of [Ni(L1)2(Cl)2]

Figure 5.21. UV-visible spectrum of [Ni(L1)2(NO3)2]

Nickel(II), 3d8

175


Figure 5.23. UV-visible spectrum of [Ni(L2)2(CH3COO)2]

Nickel(II), 3d8

176

Figure 5.24. UV-visible spectrum of [Ni(L2)2SO4]


Nickel(II), 3d8

177



Nickel(II), 3d8

178

Figure 5.28. UV-visible spectrum of [Ni(L4)2(NO3)2]


Nickel(II), 3d8

179



Nickel(II), 3d8

180


Nickel(II), 3d8

181

Table 5.4. Magnetic moment, electronic spectral data and ligand field parameters of the Ni(II) Complexes

Complexes µeff.

(BM) λmax (cm-1) Dq B β LFSE (Kjmol-1)

[Ni(L1)2Cl2] 2.97 9597, 11086, 20704, 26110 960 200 0.19 137.755 [Ni(L1)2(NO3)2] 3.01 9579, 10965, 20576, 30675 958 187 0.18 137.491 [Ni(L1)2SO4] 2.95 9671, 10989, 22936, 31645 967 327 0.31 138.821 [Ni(L1)2(CH3COO)2] 2.99 9578, 10964, 23255, 30581 958 366 0.35 137.491 [Ni(L2)2Cl2] 2.98 9506, 11050, 20576, 34602 951 207 0.2 136.446 [Ni(L2)2(NO3)2] 3.02 9488, 10787, 23474, 44643 949 387 0.37 136.187 [Ni(L2)2SO4] 3.01 9542, 10799, 21930, 37453 954 274 0.26 136.966 [Ni(L2)2(CH3COO)2] 2.96 9653, 11123, 21459, 37037 965 242 0.23 138.553 [Ni(L3)2Cl2] 2.98 9479, 12642, 27322, 35714 948 769 0.74 136.058 [Ni(L3)2(NO3)2] 3.03 9398, 12092, 25907, 940 654 0.63 134.907 [Ni(L3)2SO4] 2.97 9268, 12285, 18622, 34843 927 207 0.2 133.031 [Ni(L3)2(CH3COO)2] 3.01 9320, 11261, 20202 932 234 0.22 133.775 [Ni(L4)2Cl2] 2.95 9416, 10152, 27322 942 615 0.59 135.161 [Ni(L4)2(NO3)2] 3.03 9625, 17361, 26316 962 987 0.95 138.153 [Ni(L4)2SO4] 2.98 9149, 11377, 27933 915 791 0.76 131.327 [Ni(L4)2(CH3COO)2] 2.96 9141, 17094, 25641 914 1021 0.98 131.207 [Ni(L5)2Cl2] 2.94 9699, 13141, 27322 970 758 0.73 139.225 [Ni(L5)2SO4] 2.99 9597, 11377, 27933 960 701 0.67 137.755 [Ni(L6)2Cl2] 2.95 9328, 13441, 27174 933 842 0.81 133.9 [Ni(L6)2SO4] 3.01 9294, 13055, 27174 929 823 0.79 133.402 [Ni(L6)2(CH3COO)2] 3.02 10173, 11186, 30211, 37175 1017 725 0.7 146.023

5.5. CONCLUSION Thus on the basis of elemental analyses, molar conductance measurements, magnetic moment susceptibility, FT-IR, electronic and EPR spectral studies the following structures may be proposed for Ni(II) complexes [Figures 5.33-5.43].

Nickel(II), 3d8

182

CCH3

N

O

HN

C

H2NN

O

NH

C

NH2

CH3C

Br

Ni

X

Br

X

(Where X = Cl-, NO3-, CH3COO-)

Figure 5.33. Ni(II) complexes with ligand L1

O

OO

S

O

CH2N

NH

N

H3C

Br

O

H2NN

Ni

H3C

Br

HN

O

Figure 5.34. Ni(II) sulphate complex with ligand L1

Nickel(II), 3d8

183

CCH3

N

S

HN

C

H2NN

S

NH

C

NH2

CH3C

Br

Ni

X

Br

X



H3C

N

HN

CH2N

S

Br

H3C

N

NH

CH2N

S

Br

O

O

SNi

O

O


Nickel(II), 3d8

184

NHN

H2N O N NH

NH2ONi

X

X



O

NH2N C

O

NH

O

O O

S

ONi

NNH

CH2N


Nickel(II), 3d8

185

NHN

H2N S N NH

NH2SNi

X

X

(Where X = Cl-, NO3-, CH3COO-) Figure 5.39. Ni(II) complexes with ligand L4

S

NH2N C

S

NH

O

O O

SO

Ni

NNH

CH2N


Nickel(II), 3d8

186

O

NHN

H2N O N NH

NH2O

O

Ni

X

X

Figure 5.41. Structure of complex [Ni(L5)2Cl2]

O

H2N C

O

NH

O

O O

S

ONi

O

NNH

CH2N

O

N


Nickel(II), 3d8

187

O

NHN

H2N S N NH

NH2S

O

Ni

X

X

(Where X = Cl-, CH3COO-)


S

H2N C

S

NH

O

O O

S

ONi

O

NNH

CH2N

O

N


Nickel(II), 3d8

188

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