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Thiosemicarbazone are very versatile ligands. They can coordinate to metals as neutral molecules or after deprotonation, as anionic ligands, and can adopt a variety of different coordination modes. These derivatives are widely used in medicine for treating various diseases. [1–3] All of them possess a broad set of donor atoms and may react with metal ions to form coordination compounds of varied structure and composition. [4–6] It is established [7–10] that in many cases the biological activity of the above-mentioned drugs correlates with their ability to complex formation. Little class of coordination compounds has been subjected to as much attention as amino-complexes [11] formed by the amines with carbonyl derivatives. They constitute an interested family of ligands for bioinorganic chemistry purposes, mainly for medicinal inorganic chemistry research. Thiosemicarbazones are of considerable interest because of their chemistry and potentially beneficial biological activities, [12] such as antitumor,[13] antibacterial,[14, 15] antiviral [16] and antimalarial activities.[17, 18]

The biological activities of thiosemicarbazones are considered to be due to their ability to form chelates with metals.[19] The biological activities of such ligands are related to its chelating ability with transition metal ions, bonding through nitrogen-nitrogen-sulfur or oxygen-nitrogen-sulfur atoms. The presence of hard O- and N- and soft S-donor atoms in the backbones of these ligands enable them to react readily with both transition group and main group metal ions, yielding stable metal complexes, some of which have been shown to exhibit interesting physico-chemical properties and significant biological activities.[20-25] Transition metal complexes are powerful catalysts for organic reactions when suitable ligands are associated with the metal center, they can offer chemio, regio, or stereo selectivity under mild conditions. In view of above application, the present work relates to the Synthesis and physicochemical studies on bidentate copper (II), nickel (II) and cobalt (II) metal complexes of (E)-2-((3-hydroxy-1H-inden-1-yl)methylene)-N-((E)-prop-1-en-1-yl)hydrazine-1-carbothioamide.

1.1           Synthesis of the ligand
In a round bottom flask, a hot methanolic solution (25 ml) of 3-hydroxy-1H-indene-1-carbaldehyde (1.25 mL, 0.01 m mol) was added to a 5% acetic acid (0.32 g, 0.01 mmol) in water solution (20 mL) of (E)-N-(prop-1-en-1-yl)hydrazinecarbothioamide (1.31g, 0.01 m mol) methanol (15mL) and the reaction mixture was refluxed for about 2h. Still over the two nights at 4 ⁰C, yellowish product was obtained. After that the product was re-crystallized in the hot methanol and washed several times with diethyl ether and then evaporates the solvent under reduced pressure to afford the product. (Scheme 1)

Scheme 1 Synthesis of the 2,4-[(allyl)iminobenzyl] thiosemicarbazone
1.1           Synthesis of cobalt and nickel metal complex
An ethanolic solution of ligand (0.002 m mol, 0.05g) was added drop wise to a 0.001 m mol methanolic solution of metal chloride (cobalt and nickel) salts and refluxed at 55 ⁰C for 6 h. One-two nights standing; colored compounds are obtained as shown as in scheme 2,  the reaction mixture was monitored by thin layer chromatography. The organic layers were collected and were taken in hot methanol and washed with diethyl ether and then evaporate the solvent under reduced pressure to afford the product.

Scheme 2 Route for the synthesis of the metal complexes
1.1           Synthesis of copper complex
A quantity of (0.002 m mol, 0.05 g) of ligand was dissolved in 100 mL methanol and a solution of copper chloride (0.001 m mol, 0.017 g) in 25 mL methanol was added drop wise with continuous stirring about 12h; dark bluish single crystals appeared after one night standing (scheme 3). The resulting precipitates were filtered and washed with acetone, methanol and dried over anhydrous calcium chloride in desicator.

Scheme 3 Route for the synthesis of the copper complex

Analytical data of ligand and their metal complex

Yield: 60%; M.P. 240 ⁰C, Mol. wt. 273, color: yellow; analytical data for C₁₄H₁₅N₃OS found (calc.): C, 61.70 (61.99); H, 5.22 (5.91); N, 15.42 (15.97). IR (KBr): cm^-1 1637 n_C-NH, 1690 n_C=O, 1298 n_NH-NH. ESI-MS, m/z Data found (calc.): 241 (240), ¹H NMR (DMSO-d₆) d ppm: 7.1 (m, 8H, HC-Ar), 3.8 (s, 2H, NH-NH). ¹³C NMR (DMSO-d₆) d ppm: 117.53-121.07 (10C, CH-Ar.), 143.23 (1C, C-N), 153.88 (2C, C=O).

Cobalt complex

Yield: 28%; M.P.: 285˚C; Mol. wt. 789; color: dirty brownish; analytical data for [Co(C₂₈H₃₀N₆O₂)Cl₂] found (calc.): C, 42.58 (42.15); H, 3. 80 (3.55); N, 10.64 (10.47); IR (KBr cm^-1) 3412 n _NH, 1690 n_C=O, 3015 n_C-H, 2220 n_C-N, 519 n _M-N, 380 n _M-Cl; ¹H NMR (DMSO-d₆) d ppm: 7.3 (m, 16H, HC-Ar), 3.2 (s, 4H, NH-NH). ¹³C NMR (DMSO-d₆) d ppm: 117.53-121.07 (16C, CH-Ar.), 140.20 (4C, C-N), 153.21 (4C, C=O).

Nickel complex

Yield: 38%; M.P.: 260 ˚C; Mol. wt. 720; color: dark greenish; analytical data for [Ni(C₂₈H₃₀N₆O₂)]Cl₂ found (calc.): C, 46.66 (46.15), H, 4.16 (4.11), N, 11.66 (11.47). IR (KBr, cm^-1) 3402 n _NH, 1690 n_C=O, 3015 n_C-H, 2210 n_C-N, 511 n _M-N, 395 n _M-Cl. ¹H NMR (DMSO-d₆) d ppm: 7.1 (m, 16H, HC-Ar), 3.8 (s, 4H, NH-NH). ¹³C NMR (DMSO-d₆) d ppm: 117.53-121.07 (16C, CH-Ar.), 143.23 (4C, C-N), 153.88 (4C, C=O).

Copper complex

Yield: 38%; M.P.: 260 ˚C; Mol. wt. 720; color: dark bluish; analytical data for [Cu(C₂₈H₃₀N₆O₂)]Cl₂ found (calc.): C, 46.66 (46.15), H, 4.16 (4.11), N, 11.66 (11.47). IR (KBr, cm^-1) 3402 n _NH, 1690 n_C=O, 3015 n_C-H, 2210 n_C-N, 511 n _M-N, 395 n _M-Cl. ¹H NMR (DMSO-d₆) d ppm: 7.1 (m, 16H, HC-Ar), 3.8 (s, 4H, NH-NH). ¹³C NMR (DMSO-d₆) d ppm: 117.53-121.07 (16C, CH-Ar.), 143.23 (4C, C-N), 153.88 (4C, C=O).

The analytical data indicate that all the nickel and cobalt complexes have the general composition [Co/Ni(MIBTS)₂ Cl₂] and [Cu(MIBTS)₂] Cl₂. All the complexes are stable in air and soluble in methanol, ethanol chloroform, and DMSO/ DMF solution at room temperature. The value of molar conductance of complexes in DMSO indicates that the [Cu(MIBTS)₂]Cl₂ complexes are non electrolytes and [Cu(MIBTS)₂]Cl₂ are electrolyte. Magnetic moments lie in the range 5.01-5.08 B.M., 2.82-2.93 B.M. and 1.82-1.91 B.M. for Co(II), Ni(II) and Cu(II) complexes, respectively. The Infrared spectra of the complexes were obtained in the range 4000- 400cm^-1 region (as shown as fig. 1). Several important observations concerning the mode of coordination in these complexes are possible from these data. These are: (a) the carbonyl stretching frequencies provide a very important clue in the elucidation of structures. In a unionized carboxylate group, the C=O stretching vibration appears at 1680-1700 cm^-1. It is a well known fact that when (i) Infrared spectra of a large number of amine and ammine complexes have been examined by several authors, generally with satisfactory agreement as to exact frequencies in a given compound. The assignment of absorption in the 3000-3300 cm^-1 region to N-H stretching frequencies is beyond doubt. It is apparent from the available data that coordination of amines lowers the N-H stretching frequencies by 100-150 cm^-1. It has been proposed that the major cause of this lowering is the drainage of electronic from the nitrogen atom which, in turn, weakens the N-H bond. In the absence of these data on some of the free ligands or their salts we compare the spectral band especially above 3000 cm^-1 in the metal complexes with their corresponding ligands. Almost all the absorption bands observed above 3000 cm^-1, assigned to N-H or NH₂ stretching vibrations, are shifted to lower wave numbers in the complexes as compared to the corresponding free ligands. We interpret the observed shifts to lower energy in the complexes as arising from nitrogen coordination to the metal ions. (ii) N-N stretching vibration can be assigned on the basis of preceding work. It has been shown that if the ligands radical is attached with a conjugated system then N-N stretching occurs at about 1000 cm^-1 , either if the resulting group is chelated or not. The distinctions between chelated and non-chelated ligands have been made only by broadening of this band in the chelates. In all our complexes ν (N-N) vibration appears in the 1030-950 cm^-1 region. The slightly broad character of this band in all complexes may be taken as evidence for the nitrogen involvement in the bond formation. The IR spectra of the free ligand shows medium band at 3,253 cm^–1 due to υ (NHR) vibrational modes. This band very similar in the spectra of complexes suggested non participating in the coordination. A strong bands observed at 1,177 – 1,196 cm^–1are assigned to υ (C=S) band and a new band is observed at 332 – 336 cm^–1, indicating the bond formation between metal ions and sulphur atom.^[26] The strong band is observed at 1,544 – 1,590 cm^–1 due to υ (>C=N-) in the free ligand is shifted to lower frequency in the complexes suggested the involvement of azomethine nitrogen in chelation. The non-ligand bands at 521 – 578, 322 – 336 and 255 – 278 cm^–1 are tentatively assigned to υ (M – N),[26] υ (M – S)[27] and υ (M – Cl),[28] respectively. [26] ¹H NMR spectra of the ligand, the signals of the –NH protons were observed as singlets at δ 11.42 –11.83. These signals also appeared in the ¹H NMR spectra of the metal complexes indicating non coordinating the metal ions. The signals of the –CH=N proton which appears as singlets at δ 8.03 – 8.17 in the ligand show a shift to downfield in δ 0.03 – 0.08 after complexation. The shift indicates the coordination of the imine nitrogen [29] to the metal center. The signals of the aromatic protons of the ligand appeared at δ 7.21 – 7.91, and the resonance lines found correspond to the calculated multiplicity. These signals do not suffer relevant changes in the chemicals shifts for the metal complexes. The signals of the –C₃H₅ proton appears at δ 2.34 – 2.48 in the spectrum of ligand and complexes. The –CH=CH protons signals appears at δ 1.02 – 1.09 in the ligand and metal complexes spectra.  The Ni(II) and Co(II) complexes are non-electrolytic and , Cu(II) complex electrolyte in nature by their molar conductivity (Lm) as measured in DMSO in the range 96–104 Ω^–1 cm² mol^–1 and  93 Ω^–1 cm² mol^–1, respectively.[30] Mass spectra provide a vital clue for elucidating the structure of compounds. The spectrum shows the molecular ion peak at m/z = 245 (C₁₃H₁₅N₃S, calculated atomic mass 244 amu due to ¹³C and ¹⁵N isotopes) and other peaks like 44, 60, 78, 88, 91, 119, and 177 may be due to different fragments. The different competitive fragmentation pathways of ligand give the peaks at different mass numbers at 245. The intensity of these peaks reflects the stability and abundance of the ions. The weak peak described at 135 amu is assigned to the fragment [C₆H₈N₄] ⁺, corresponding to the loss of CS group. A very weak peak at 119 amu is assigned to the fragment [C₆H₆N₃] ⁺, corresponding to the loss of CSNH₂ group. The most intense peak at 91 corresponding to the fragment [C₆H₅N] ⁺. Other peak at 88, 78, 60, and 44 correspond to fragments [CH₃N₃S] ⁺, [C₅H₄N] ⁺, [CSNH₂] ⁺, and [CS] ⁺, respectively. The intensity of theses peaks gives an idea of the stability of these fragments. The EPR spectra of the Cu (II) complexes were recorded as polycrystalline sample at LNT since the rapid spin lattic relaxtion of Cu (II) broaden the lines at higher temperature. The g_II value for metal complexes is less than 2.3 suggesting a small amount of ionic character of the metal-ligand bond. The trend g_II > g_^ >2.0023, suggests that the unpair electron lie predominantly in the dx²–y² orbital characteristic of octahedral geometry [31] in Cu (II) complexes. The electronic spectra of the complexes recorded in MeOH solution are given table 1. The electronic spectra of the cobalt(II) complex showed three bands at 8780-8810, 17475-17775 and 30235-30270 cm^-1, which may be assigned to ⁴T_{1g →} ⁴T_2g (F), ⁴T_1g → ⁴T_1g(P), and ⁴T_{1g →} ³A_2g (F) transitions and suggested octahedral geometry around the cobalt ion.[32] The cobalt complexes showed magnetic moment values 4.70-490 B.M. at room temperature. These high values of the magnetic moments and the stoichiometries suggest a coordination number of six for the central cobalt ions and an octahedral geometry. The magnetic properties of copper complex may be divided into broad classes. First, are those having essentially temperature independent magnetic moments in rang 2.20 BM. Those exhibiting such moments are mononuclear complexes having no major interaction between the unpaired electrons on different copper ion. The moments in such complex, as in apparent, lie appreciably above the spine-only value (1.73 BM), but as the electronic ground states are non-degenerate this cannot arise from inherent angular momentum in the ground state. It arises due to mixing in of some orbital angular momentum from excited states via spin orbit coupling. The copper (II) complexes exhibit magnetic moments of 1.70-1.75 B.M., respectively, at room temperature. These values are quite close to the spin-allowed values expected for an S= 1/2 system and may be indicative of a square planner geometries around copper (II) ions. The electronic spectra of the copper(II) complexes display a broad band at 14220-14918 cm^-1 due to ²B_1g → ²E_g and two bands at 16390-16550 and 27250-27350 cm^-1 assigned to d–d transitions and a charge transfer band respectively, of square planner environment.[33] The observed magnetic moment of the copper complexes are 1.75 -180 BM. The nickel (II) complexes exhibited three bands 9960-10565, 15850-17155, and 24200-29985 cm^-1 assignable to the transitions ³A_2g (F) → ³T_2g (F) (ν₁), ³A_2g (F) → ³T_1g (F) (ν₂) and ³A_2g (F) → ³T_1g (P) (ν₃), respectively which are characteristic of nickel (II) in octahedral geometry.[34] The magnetic moment of the nickel complexes are 3.20-3.25BM, were lay in the range of octahedral geometries.

 

Table 1 Electronic spectral data (nm) of the metal complexes

Complexes	λmax  (nm)	Assignments
[Co (MIBTS)2 Cl2]	211,283, 333, 399, 457, 643	n–p*, p–p*, d–d
[Cu (MIBTS)2 Cl2]	236, 275, 335, 366, 425, 580	n–p*, p–p*, d–d
[Ni (MIBTS)2 Cl2]	270, 330, 350, 375, 610	n–p*, p–p*, d–d

 

On the basis of the above observations, it is tentatively suggested that all of the complexes show an octahedral geometry in which the two ligand act as bidentates. These possibly accommodate them selves around the metal atom in such a way that a stable chelate ring is formed giving, in turn, stability to the formed metal complexes.

Microbiology Essay

For the antibacterial and antifungal assays, the compounds were dissolved in dimethylformamide. Further dilutions of the compounds and standard drugs in the test medium were prepared at the required quantities of 500 and 1000 ppm concentrations with dextrose broth. The minimum inhibitory concentrations (MIC) were determined using the two fold serial dilution technique.^[35] A control test was also performed containing inoculated broth supplemented at the same dilutions used in our experiments and found inactive in the culture medium. All the compounds were tested for their in vitro growth inhibitory activity against different bacteria and the fungus. Origins of bacterial strains are S. aureus ATCC 29253, S. aureus ATCC 3160, as Gram-positive. Gentamycin and Amphotericin B were used as control drugs. The data on the antimicrobial activity of the compounds and the control drugs as MIC values are given in table 2.The cultures were obtained from SRL broth for all the bacterial strains after 24 h of incubation at 37 ⁰C. C. albicans were maintained in dextrose broth after incubation for 24 h at 25 ⁰C. Testing was carried out in dextrose broth at pH 7.4 and the two fold serial dilution technique was applied. A set of tubes containing only inoculated broth was used as controls. For the antibacterial assay after incubation for 24 h at 37 ⁰C and after incubation for 48 h at 25 ⁰C for the antifungal assay, the last tube with no growth of microorganism and/or yeast was recorded to represent the MIC expressed in ppm. Every experiment in the antibacterial and antifungal assays was replicated twice and the data is given in table 2. The observation on the biological assay indicate that the antibacterial action due to all compounds have NOS group which is of considerable chemotherapeutic interest. From the table 2 fig. 2 (a), (b)and bar graphs it is evidence that among all the newly synthesize compounds tested for their antibacterial and antifungal activities against Staphylococcus aureus (ATCC 25923), Staphylococcus aureus (ATCC 3160), Cabdida albicans (227) and Staphylococcus cereviscae (361) were determined as MIC values.

All the investigated compounds showed good activity against S. aureus. The zone of inhibition in ml of the test compounds against the micro-organism Staphylococcus aureus (ATCC 25923), Staphylococcus aureus (ATCC 3160), Cabdida albicans (227) and Staphylococcus cereviscae (361). the data indicate that among the bacterias employed, S. aureus is found to be more sensitive to these compounds where as the gram negative bacteria, E. coli shows resistance to most of the compounds. [34, 35] The zone of inhibition tabulated reveal that the antibacterial activity of the compounds is specific to the microorganism examined. Analysis of the data showed that the general the fungi C. albicans (227) was more susceptible to the irreversible toxic effects of screened compounds than S. cereviscae (361). Variation in the response of fungi studies to chemical screened may be attributed to the tolerance of them by test fungi. The effects of assayed chemicals on test fungi differed in accordance with the concentrations used. Generally fungi toxicity enhanced with the increase in the dose of compounds. Higher the concentration longer was the persistence of chemicals.

Figure 3 Showing the bar graph antibacterial (a) and antifungal activities (b) at 1000 ppm concentration after 48

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