Spectroscopic, cyclic voltammetric and biological studies of transition metal complexes with mixed nitrogen–sulphur (NS) donor macrocyclic ligand derived from thiosemicarbazide

Spectrochimica Acta Part A 62 (2005) 453–460

Spectroscopic, cyclic voltammetric and biological studies of transition metal complexes with mixed nitrogen–sulphur (NS) donor macrocyclic ligand derived from thiosemicarbazide Sulekh Chandra ∗ , Lokesh Kumar Gupta, Sangeetika Department of Chemistry, Zakir Husain College, University of Delhi, J.L. Nehru Marg, New Delhi 110002, India Received 25 October 2004; received in revised form 24 January 2005; accepted 24 January 2005

Abstract The complexation of new mixed thia-aza-oxa macrocycle viz., 2,12-dithio-5,9,14,18-tetraoxo-7,16-dithia-1,3,4,10,11,13-hexaazacyclooctadecane containing thiosemicarba-zone unit with a series of transition metals Co(II), Ni(II) and Cu(II) has been investigated, by different spectroscopic techniques. The structural features of the ligand have been studied by EI-mass, 1 H NMR and IR spectral techniques. Elemental analyses, magnetic moment susceptibility, molar conductance, IR, electronic, and EPR spectral studies characterized the complexes. Electronic absorption and IR spectra of the complexes indicate octahedral geometry for chloro, nitrato, thiocyanato or acetato complexes. The dimeric and neutral nature of the sulphato complexes are confirmed from magnetic susceptibility and low conductance values. Electronic spectra suggests square-planar geometry for all sulphato complexes. The redox behaviour was studied by cyclic voltammetry, show metal-centered reduction processes for all complexes. The complexes of copper show both oxidation and reduction process. The redox potentials depend on the conformation of central atom in the macrocyclic complexes. Newly synthesized macrocyclic ligand and its transition metal complexes show markedly growth inhibitory activity against pathogenic bacterias and plant pathogenic fungi under study. Most of the complexes have higher activity than that of the metal free ligand. © 2005 Elsevier B.V. All rights reserved. Keywords: Thiosemicarbazide; Complexes; N–S donor

1. Introduction The design and synthesis of macrocyclic ligand having heteronucleating donor atoms has been increased significant interest. It offers exciting possibilities for creative minds to construct novel supramolecular assemblies that are capable of performing highly specific molecular functions [1]. The N–S donor macrocycles also have theoretical interest, as they are capable of furnishing an environment of controlled geometry and ligand field strength [2]. The precise molecular recognition between macrocyclic ligands and their guest provides a good opportunity for studying key aspects of supramolecular chemistry, which are also significant in a variety of disciplines including chemistry, biology, physics, medicine and related ∗

Corresponding author. Tel.: +91 11 22911267; fax: +91 11 23215906. E-mail addresses: schandra [email protected] (S. Chandra), lokesh [email protected] (L.K. Gupta). 1386-1425/$ – see front matter © 2005 Elsevier B.V. All rights reserved. doi:10.1016/j.saa.2005.01.015

science and technology [3]. Chemically mixed donor macrocycles are important because of great versality as ligands due to presence of several potential donor atoms, their flexibility and ability to coordinate with several metal ions [4]. Biological activities of the complexes may be related to the redox properties of complexes. Some copper complexes exhibit the abilities of superoxide dismutase (SOD) and chemical nucleases [5,6]. The recent growing interest in the electrochemistry (reduction and oxidation process) of macrocyclic complexes derived from recognition of biological importance of the less common oxidation states of Cu and Ni [7,8]. Redox potentials of Cu(II)/Cu(I) depends on the relative thermodynamic stabilities of the two oxidation states in a given ligand environment. The structural features including ring size, degree and arrangement of unsaturation and substitution [9,10]. In the present paper we report the synthesis and characterization of a series of cobalt(II), nickel(II) and copper(II) complexes obtained from 2,12-dithio-5,9,14,18-tetraoxo-7,16-

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dithia-1,3,4,10,11,13-hexaazacyclooctadecane. In addition we have studied the electrochemical behaviour of complexes by cyclic voltammetry.

2. Experimental All the chemicals used were of AnalaR grade and were procured from, Sigma–Aldrich, E-merck and Fluka. 3. Preparation of ligand A warm ethanolic (50 mL) solution of thiosemicarbazide (0.01 mol., 0.92 g) was added to a warm ethanolic (50 mL) solution of thiodiglycolic acid (0.01 mol., 1.50 g). After addition the mixture was refluxed for 7–8 h. On keeping it upto 5–10 ◦ C, overnight, a white coloured crystalline solid was formed, which was filtered, washed with cold ethanol and dried in vacuo over P4 O10 . Yield 65%, m.p. 165 ◦ C (molecular weight found 346: C, 34.23; H, 4.45; N, 24.03; calculated 10 H14 N6 S4 : C, 34.66%; H, 4.07%; N, 24.25%). 4. Preparation of complexes

solution on a Shimadzu UV mini-1240 spectrophotometer. Molar conductance is measured on an ELICO Conductivity Bridge (Type C M 82 T). Magnetic susceptibility measurements were made on Gouy Balance at room temperature using CuSO4 ·5H2 O as calibrant. Electron Impact Mass spectra were recorded on JEOL, JMS, DX-303 mass spectrometer. 1 H NMR spectrum was recorded on Hitachi FT-NMR, Model R-600 spectrometer using CDCl3 as solvent. Chemical shifts are given in ppm relative to tetramethylsilane. EPR spectra of the complexes were recorded as powder samples at room temperature (RT) and liquid nitrogen temperature (LNT) on an E-4 EPR spectrometer using DPPH as the g-marker. Molecular weight of the complexes was determined in benzene (freezing point). Electrochemical properties were performed using platinum wire as auxiliary, Ag/AgCl as reference electrode and glassy carbon as working electrode at RT. Cyclic voltammetry studies of complexes were carried out in 0.01 M solutions in dimethylsulfoxide containing [NBu4 ][PF6 ] as supporting electrolyte. The range of potential was studied in between +1.5 and −2.0 V. Before ploting graph nitrogen gas was passed in each solutions for 5 min.

6. Results and discussion

A warm ethanolic (10 mL) solution of corresponding metal salt (0.001 mol.) was added to a warm ethanolic (10 mL) suspension of the macrocyclic ligand (0.001 mol., 0.34 g). The mixture was heated under reflux with stirring for 6–7 h. On cooling a coloured complex precipitated out, which was filtered washed with cold ethanol and dried in vacuo over P4 O10 (yield 50–62%). 5. Physical measurements Microanalysis (C, H and N) of these complexes were carried out on a Carlo-Erba 1106 elemental analyzer. IR spectra were recorded on FTIR Spectrum BX-II Spectrophotometer KBr pellets. Electronic spectra were recorded in DMSO

Physical characterization, microanalytical and molar conductance data of the complexes are given in Table 1. The analytical data of the complexes correspond well with the general formula MLX2 and M2 L(SO4 )2 where L = macrocyclic ligand; M = Co(II), Ni(II) or Cu(II) and X = Cl− , NO3 − , SCN− or CH3 COO− . The molar conductance indicates that the all the complexes are 1:2 electrolyte in nature except sulphato complexes which are non-electrolytic in nature. The EI mass spectrum of L (Fig. 1) confirms the proposed formula by showing a peak at 345 amu (M ± 1) corresponding to the molecular ion [C10 H14 N6 S4 ]+ . It also shows peak at 254 amu corresponds to M–N3 H5 CS+ (Fig. 2). These data suggests the (2 + 2) complete condensation of thiodiglycolic acid with thiosemicarbazide.

Fig. 1. Synthesis and structure of ligand (L).

N H

2.32 (2.61) 2.10 (2.37) 1.65 (1.95) 2.67 (2.41) 2.34 (2.61) 2.56 (2.37) 1.59 (1.95) 2.15 (2.41) 2.12 (2.58) 2.56 (2.35) 1.64 (1.93) 3.18 (3.41)

C

22.78 (22.22) 20.74 (20.23) 16.98 (16.67) 24.17 (24.61) 22.56 (22.23) 20.57 (20.24) 16.43 (16.68) 24.38 (24.62) 22.19 (22.03) 19.87 (20.08) 16.21 (16.45) 28.10 (28.39)

NMR spectral data of L in CDCl3 show the signals corresponding to proposed structure, as it does not show any signal corresponding to primary amine group and alcoholic protons. It gives multiplet in the region of δ9.21–8.96 ppm of amide protons (CO–NH)(4H) and (CS–NH)(2H). It also gives a singlet in the region of δ2.54–3.25 ppm assignable to methylene protons (CO–CH2 –S)(8H). The infrared spectrum of the ligand does not exhibit any band corresponding to free –NH and –OH groups, and the appearance of four new bands characteristics of amide groups at 1643 ν(C O), amide I, 1620 ν(CS–NH), amide II, 1531 ν(C–N) + δ(N–H), amide III, 1287 δ(N–H) and 647 cm−1 φ(C O), amide IV (Table 2), which support to macrocyclic nature of the ligand. A sharp band observed in the region 3372–3179 cm−1 may be assigned to ν(N–H) of the secondary amino group [11]. The proposed structure of the ligand is given in Fig. 2. In the spectra of complexes the appearance of a new band at 380–450 cm−1 which may be assigned to ν(M–N). It supports the involvement of nitrogen in coordination. A strong to medium intensity band in the region of 750–780 cm−1 has been assigned to ν(C–S). The ν(C–S) band is shifted by ca. 30 cm−1 in the complexes clearly indicate that sulphur is also takes place in coordination.

11.57 (11.84) 9.54 (9.93) 16.12 (16.35) 10.39 (10.06) 10.65 (10.86) 9.45 (9.89) 16.12 (16.30) 10.45 (10.02) 11.39 (11.66) 10.52 (10.62) 17.19 (17.41) 10.57 (10.73)

M

1H

210 205 250 229 210 205 250 229 210 205 250 229 265 250 10.0 230 215 220 10.0 225 265 250 10.0 195

533 (540.36) 590 (593.46) 727 (720.51) 580 (585.62) 533 (540.12) 6005 (93.22) 714 (720.04) 579 (585.37) 539 (544.97) 590 (598.07) 720 (729.74) 587 (592.15)

Found (calculated) (%)

m.p.a (◦ C) Molar condition (
−1 cm2 mol−1 )

Molar weight found (calculated)

Fig. 2. Electron impact mass spectrum of ligand (L).

Decomposition temperature.

Green Green Red Light brown Green Light purple Light green Light green Royal blue Greenish blue Blue Marine blue [CoH6 L]Cl2 CoC10 H14 N6 S4 O4 Cl2 [CoH6 L](NO3 )2 CoC10 H14 N8 S4 O10 [Co2 H6 L(SO4 )2 ]Co2 C10 H14 N6 S6 O12 [CoH6 L](SCN)2 CoC12 H14 N8 S6 O4 [NiH6 L]Cl2 NiC10 H14 N6 S4 O4 Cl2 [NiH6 L](NO3 )2 NiC10 H14 N8 S4 O10 [Ni2 H6 L(SO4 )2 ]Ni2 C10 H14 N6 S6 O12 [NiH6 L](SCN)2 NiC12 H14 N8 S6 O4 [CuH6 L]Cl2 CuC10 H14 N6 S4 O4 Cl2 [CuH6 L](NO3 )2 CuC10 H14 N8 S4 O10 [Cu2 H6 L(SO4 )2 ]Cu2 C10 H14 N6 S6 O12 [CuH6 L](CH3 COO)2 CuC14 H20 N6 S4 O8

a

Colour

6.1. Bands due to anions

Complex molecular formula

Table 1 Physical, analytical and molar conductance data of the complexes

455

15.21 (15.55) 18.39 (18.88) 11.39 (11.66) 19.67 (19.13) 15.23 (15.55) 18.47 (18.88) 11.54 (11.67) 19.4519.14) 15.10 (15.42) 18.38 (18.73) 11.03 (11.51) 14.38 (14.19)

S. Chandra et al. / Spectrochimica Acta Part A 62 (2005) 453–460

The IR spectra of nitrato complexes indicate that the nitrate group is uncoordinated as the strong band of free nitrate is present in the region of 1380–1395 cm−1 . Thus, the nitrate group is not coordinated and it is present outside the sphere in all complexes [11]. The IR spectra of sulphato complexes of Co(II), Ni(II) and Cu(II) complexes the ν3 and ν4 split in the region of 1099–1050 (3 s) (ν3 ) and 641–571 cm−1 (ν4 ) into three bands. These results suggest that, the symmetry is further lowered and reduced to C2v and indicates bidentate nature of sulphate group [11]. Thus sulphate group acts as bridging bidentate group. The acetato complexes of copper show the IR band at 1589 and 1444 cm−1 corresponding to ionic acetate ion. The thiocyanato complex of Co(II) and Ni(II) gives IR peaks in the region 2105–2112 cm−1 corresponds to uncoordinated thiocyanate group [11,12] (Fig. 3).

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Table 2 Relevant IR spectral peaks (cm−1 ) and their assignment Complexes

[CoH6 L]Cl2 [CoH6 L](NO3 )2 [Co2 H6 L(SO4 )2 ] [CoH6 L](SCN)2 [NiH6 L]Cl2 [NiH6 L](NO3 )2 [Ni2 H6 L(SO4 )2 ] [NiH6 L](SCN)2 [CuH6 L]Cl2 [CuH6 L](NO3 )2 [Cu2 H6 L(SO4 )2 ] [Cu2 H6 L(CH3 COO)2 ]

ν(N–H)

3249 3143 3104 3156 3150 3151 3120 3135 3150 3182 3169 3163

Amide bands I

II

III

IV

1640 1637 1642 1642 1630 1639 1662 1654 1645 1632 1642 1596

1561 1555 1545 1530 1535 1520 1535 1530 1561 1603 1573 1560

1213 1243 1247 1254 1240 1256 1230 1246 1240 1235 1231 1232

768 822 800 785 725 700 707 700 700 719 700 716

6.2. Cobalt(II) complexes The magnetic moment of all the complexes except sulphato complex lie in the range of 4.98–5.08 B.M. corresponding to three unpaired electrons. The value of magnetic moments depends upon the amount of total orbital angular momentum and total spin angular momentum. The sulphato complexes show magnetic moment 1.98 B.M. corresponding to one unpaired electron. Electronic spectra show bands in the region of 10165–10405 (ν1 ), 14705–17650 (ν2 ) and 21038–26609 cm−1 (ν3 ) (Fig. 4a) and are listed in Table 3. These bands may be assigned to following transitions 4 T1g → 4 T2g (F), 4 T1g → 4 A2g (F) and 4 T1g → 4 T1g (P), respectively. The position of bands suggests octahedral geometry around the Co(II) ion [13,14]. The electronic spectra of the sulphato complexes show a narrow band at 9678 and a broader band at 22,180 cm−1 , which correspond to square planar geometry [14]. The various ligand field parameters were calculated for the cobalt(II) complexes. The value of Dq has been calculated from Orgel energy level diagrams using the ν3 /ν1 ratio.

Fig. 3. IR spectrum of [Co2 H6 L(SO4 )2 ], showing bidentate sulphate group.

ν(M–N)

ν(M–Cl)

447 450 446 452 450 452 406 402 418 420 425 435

340 – – – 354 – – – 350 – – –

The nephelauxetic parameters β is readily obtained using the relation: β = B (complex)/B (free ion), where B (free ion) is 1120 cm−1 . The value of β lie in the range 0.854–0.891. The value of β indicates that the covalent character of metal ligand σ bond is low. The ligand field stabilization energy is also calculated and given in Table 4.

Fig. 4. Electronic spectra of the complexes. (a) [CoH6 L]Cl2 in DMF solution, (b) [NiH6 L]Cl2 in DMF solution, (c) [CuH6 L]Cl2 in DMF solution.

S. Chandra et al. / Spectrochimica Acta Part A 62 (2005) 453–460 3A

Table 3 Electronic spectral and magnetic moment data of the complexe Complexes

Spectral bands (cm−1 )

µeff (B.M.)

[CoH6 L]Cl2 [CoH6 L](NO3 )2 [Co2 H6 L(SO4 )2 ] [CoH6 L](SCN)2 [NiH6 L]Cl2 [NiH6 L](NO3 )2 [Ni2 H6 L(SO4 )2 ] [NiH6 L](SCN)2 [CuH6 L]Cl2 [CuH6 L](NO3 )2 [Cu2 H6 L(SO4 )2 ] [CuH6 L](CH3 COO)2

10405, 14705, 21038 10373, 17650, 23450 9678, 22180 10165, 15986, 26609 10384, 14700, 16638, 20890 10050, 16667, 21500 10030, 16666, 22000 10384, 16920, 21200 14386, 18720, 20833 13661, 18600 14210, 18650, 23200 14000, 18650, 21450

4.98 5.08 1.98 5.07 2.91 2.94 Diamagnetic 3.04 1.98 2.01 0.68 2.02

Table 4 Ligand field parameters of the complexes Complexes

Dq (cm−1 )

ν3 /ν1

B (cm−1 )

β

LFSE (kJ mol−1 )

[CoH6 L]Cl2 [CoH6 L](NO3 )2 [CoH6 L](SCN)2 [NiH6 L]Cl2 [NiH6 L](NO3 )2 [NiH6 L](SCN)2

1201 1244 1210 1038 1005 1038

2.31 2.26 2.12

975 957 998 519 534 464

0.870 0.854 0.891 0.49 0.50 0.44

114 118 115 148 144 148

The EPR spectra of Co(II) complexes were recorded as polycrystalline samples at LNT, due to the rapid spin lattice relaxation of Co(II) broadens the lines at higher temperature. The deviation of ‘g’ values from the free electron value (2.0023) may be due to angular momentum contribution in the complexes. The g values for complexes are listed in Table 5. 6.3. Nickel(II) complexes The magnetic moment lies in the range of 2.91–3.04 B.M. corresponding to two unpaired electrons. It indicates that these complexes have six coordinate octahedral geometry. The complex with molecular formula [Ni2 L (SO4 )2 ] is diamagnetic suggesting square planar geometry. Electronic spectra of [NiL]X2 complexes (where X = Cl− , NO3 − and NCS− ) show bands in the region of 10050–10384 cm−1 , 16638–16667 cm−1 and 20890–21500 cm−1 (Fig. 4b). These bands may be assigned to 3 A2g (F) → 3 T2g (F) (ν1 ), 3 A2g (F) → 3 T1g (F) (ν2 ) and Table 5 EPR spectral data of the complexes as polycrystalline sample Complexes

Temperature

g

g⊥

giso

G

[CoH6 L]Cl2 [CoH6 L](NO3 )2 [Co2 H6 L(SO4 )2 ] [CoH6 L](SCN)2 [CuH6 L]Cl2 [CuH6 L](NO3 )2 [CuH6 L](CH3 COO)2 [Cu2 H6 L(SO4 )2 ]

LNT LNT LNT LNT RT RT RT RT

4.30 4.28 4.31 4.23 2.150 2.198 2.198 2.093

4.30 4.28 4.31 4.23 2.045 2.098 2.094 2.030

– – – – 2.080 2.131 2.128 2.051

– – – – 3.333 2.020 2.106 3.101

giso = g + 2g⊥ /3 and G = (g − 2)/(g⊥ − 2).

2g (F) →

457

3T

1g (P) (ν3 ) transitions, respectively. It suggests octahedral geometry of Ni(II) complexes [14]. The electronic spectra of sulphato complex show bands at 10,030, 16,666 and 22,000 cm−1 assignable to following 1 A1g → 1 A2g (ν1 ), 1 A1g → 1 B2g (ν2 ) and 1 A1g → 1 Eg (ν3 ) transitions, respectively, corresponding to square planar geometry. The various ligand field parameters are calculated by Orgel diagrams and are compared [14].

6.4. Copper(II) complexes The magnetic moment of all the Cu(II) complexes recorded at room temperature lie in the range 1.98–2.02 B.M. (Table 1) corresponding to one unpaired electron which are higher than spin-only value, i.e. 1.73 BM for one unpaired electron. This reveals that these complexes are monomeric in nature and also shows the absence of metal–metal interaction along the axial positions. Whereas, the [Cu2 LSO4 )2 ] type complex show magnetic moment at 0.68 B.M. indicates the presence of metal-metal interaction between them. Electronic spectra of six-coordinate Cu(II) complexes have either D4h or C4v symmetry, and eg and t2g level of 2 D free ion term will split into B1g , A1g , B2g and Eg level, respectively. Thus, the three spin allowed transitions are expected in the visible and near IR region. But only few complexes are known [15] in which such bands are resolved either by Gaussian Analysis or single crystal polarisation studies. These bands may be assigned to following transitions, 2 B → 2 A (d 2 2 2 2 2 1g 1g x−y → dz ), B1g → B2g (dx−y → dxy ) and 2 B → 2 E (d 2 1g g x−y → dxz dyz ) in order of increasing energy. The electronic spectra of the complexes having molecular formula [CuL]X2 (where X = Cl− , NO3 − and CH3 COO− ) show two bands in the range of 13661–14386 cm−1 , 18600–18720 cm−1 and 20,833–21,450 cm−1 (Fig. 4c) assignable to 2 B1g → 2 A1g ,2 B1g → 2 B2g and 2 B1g → 2 Eg transitions, respectively. This suggests that these complexes have tetragonal geometry [14]. The electronic spectra of sulphato complexes with these ligands corresponds to square planar geometry. Tetragonal Cu(II) complexes give anisotropic EPR spectra. The anisotropic g values have been calculated using the following expressions: g = 2(1 − 4λ/E2 ) and g⊥ = 2(1 − λ/E3 ), Where λ = 823 cm−1 and E2 = 2 B1g → 2 B2g , E3 = 2 B1g → 2Eg . The stronger interaction along the ‘Z’-axis is to be accompained by an increase in the value of g . The stronger axial bonding leads to an increase in the length of the bond in the XY plane, which results in a decrease of both the in-plane covalency 2 and energy of the dx−y → dxy transition [15]. Both these factors tend to increase of the value of g . The EPR spectra of the complexes under study recorded as polycrystalline sample at room temperature. Spectra for polycrystalline sample exhibit absorptions typical for the mononuclear species with axial symmetry. A representative spectra of the copper(II) complexes as polycrystalline sample

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Fig. 5. EPR spectra of the complexes (a) [CoH6 L]Cl2 as polycrystalline sample at LNT, (b) [CuH6 L]Cl2 as polycrystalline sample at RT.

at RT is shown in (Fig. 5). The g value obtained from these spectra are presented in Table 5. In tetragonal and square planar geometry the unpaired 2 electron lies in the dx−y orbital giving 2 B1g as the ground state with g > g⊥ . The g and g⊥ values were computed from the spectrum using DPPH free radical as ‘g’ marker. The ‘g’ values and spin Halmiltonian parameters are summarized in Table 5. Kivelson and Neiman have reported the g value <2.3 for covalent character of the metal–ligand bond and >2.3 for ionic character. Applying this criterion the covalent character of the metal-ligand bond in the complexes under study can be predicted. The trend g > g⊥ > ge (2.0023) observed for these complexes show that the unpaired electron is localized 2 in dx−y orbital of the Cu(II) ions and the spectral features are characterstics of axial symmetry [16]. In addition there is exchange coupling interaction between two copper centers explained by Hathaway [17] and expressed as G = (g − 2)/(g⊥ − 2). According to Hathaway if the value of G is greater than 4, the exchange interaction is negligible, whereas when the value of G is less than 4 a considerable interaction is indicated in solid complex. The calculated G values are given in Table 5. On the basis of elemental analysis, molar conductance measurements, magnetic susceptibility, IR, electronic, EPR data and the subsequent discussion given above the following

structures may be proposed for all the complexes (Fig. 6a and b). 6.5. Electrochemical behavior The electrochemical properties of metal complexes, particularly with sulphur donor atoms have been studied in order to consider spectral and structural changes accompanying electron transfer. The cyclic voltammetric properties of Co(II) complexes was studied in the range from +1.5 to −2.0 V. The electrochemical behaviour of all the cobalt(II) complexes were same even sulphato complexes having different geometry proposed from spectroscopic data for these complexes. Cyclic voltammogram of [CoH6 L]Cl2 shows an quasi-reversible peak for the couple: cobalt(II) → cobalt(III) at Epa = 1.09 V with the direct cathodic peak for cobalt(III) → cobalt(II) at Epc = 0.89 V (Fig. 7). Cyclic scan in the region of 0–2.0 V it exhibits two peaks corresponding to cobalt(II) → cobalt(I) at Epc = −1.23 V and cobalt(I) → cobalt(0) at Epc = −1.59 V. In the anodic region it also exhibits two peaks shows their oxidation of cobalt(0) to cobalt(II) (Table 6). Voltammetric parameters are studied in the scan rate interval of 50–800 mV s−1 . The ratio between the cathodic peak current and the square root of the scan rate (Ipc /ν1/2 ) is

Fig. 6. (a) Structure of [ML]X2 , (b) [M2 L(SO4 )2 ] (where M = Co(II) or Ni(II) or Cu(II) and X = Cl− , NO3 − , SCN− or CH3 COO− ).

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Fig. 7. Cyclic voltammogram of (I) [CoH6 L]Cl2 (II) [CuH6 L]Cl2 .

approximately constant. The peak potential shows a small dependence with the scan rate. The ratio Ipa − Ipc is close to unity. From these data, it can be deduced that this redox couple is related to a quasirreversible one-electron transfer process controlled by diffusion. The electrochemical behaviour of all the Ni(II) complexes are similar in the same conditions and depends on the potential range. On scanning from 0.0 to −2.0 V, the cyclic voltammogram of [NiH6 L]Cl2 shows two waves in the cathodic scan at −1.00 and −1.58 V, corresponds to reduction of the Ni(II) ion and in the reverse scan two anodic peaks at −1.56 and 1.05 V. The redox couple are studied in the interval of 50–800 mV s−1 shows a linear variation of Ipc versus ν1/2 . It show small dependence of the potential peak with scan rate is observed [18,19], and again Ipa − Ipc ratio is close to unity. These suggests a diffusion controlled quasirreversible Ni(II)–Ni(I) process. In sulphato complex of nickel it is diffiTable 6 Electrochemical data of the complexes Compound

Epc (V)

Epa (V)

E1/2 (V)

Ep (V)

Ipc /Ipa

[CoH6 L]Cl2 [CoH6 L](NO3 )2 [Co2 H6 L(SO4 )2 ] [CoH6 L](SCN)2 [NiH6 L]Cl2 [NiH6 L](NO3 )2 [Ni2 H6 LSO4 ] [NiH6 L](SCN)2 [CuH6 L]Cl2 [CuH6 L](NO3 )2 [CuH6 L](CH3 COO)2 [Cu2 H6 L(SO4 )2 ]

0.89 0.85 0.85 0.80 0.75 0.65 0.73 0.70 0.82 0.70 0.73 0.78

1.09 1.15 1.20 1.12 1.05 1.00 1.18 1.08 1.17 1.10 1.12 1.14

0.99 1.00 1.02 0.96 0.90 0.83 0.95 0.89 0.99 0.90 0.93 0.96

0.20 0.30 0.35 0.32 0.30 0.35 0.45 0.38 0.34 0.40 0.39 0.36

1.06 0.99 1.02 1.03 1.00 1.06 0.11 0.99 0.80 0.89 0.99 0.81

Electrochemical data recored in mV, at room temperature, scan rate 100 mV s−1 , E1/2 = 1/2(Epc + Epa ), The ratio of Ipc /Ipa is constant for scan in the range of 50–800 mV s−1 .

cult to have access +1 state for nickel when ligand in neutral form. The positive values of the oxidation process show that it is easier to oxidize as compared to reduce. The electrochemical properties of the Cu(II) complexes were carried out in the range from +1.5 to −2.0 V. both the ranges of potentials were studied independently. The cyclic voltammogram between +1.5 and 0 V show oxidation process Cu(III)/Cu(II). While in the range of 0 to −2.0 V it show the reduction processes which may be assigned Cu(II)/Cu(I) and Cu(I)/Cu(O) processes (Table 6). The voltammogram of copper complexes between +1.5 and 0 V shows cathodic peak in the region of +1.25 V and in forward scan at +0.825 V. They gives anodic peak near +1.17 V associated with cathodic peak at 0.830 V Fig. 7. The Ep values lies in the range of 60–80 mV, and the scan rate in vary in the range of 50–800 mV s−1 , the ratio of between the cathodic peak current ant square root of the scan rate (Lpc /ν1/2 ) is also constant. All of them resemble of a quasireversible one-electron transfer process [20,21]. Cyclic scan in the range of 0 to −2.0 V reveals peak at −0.7.5 and −1.15 V in the cathodic scan and the forward scan shows a peak at −0.680 V associated with cathodic peak at −1.10 V. These peaks corresponds to successive Cu(II) reduction processes. The voltammogram of chloro, nitrato and acetato complexes are quite similar [22]. 6.6. Biological screening The antibacterial activity of the macrocyclic ligand and its soluble transition metal complexes was performed by the disc diffusion technique [23,24]. The zone of inhibition was measured against Staphylococcus aureus, Escherchia coli and Salmonella albus. A clearing zone around the discs indicates the inhibitory activity of the compound on the organism.

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References

Graph 1. Biological activity of the ligand and complexes.

Results are shown in Graph 1, clearly indicate that the inhibition are much larger by metal complexes as compare to the metal free ligand. The increased activity of the metal chelates can be explained on the basis of chelation theory [25]. The chelation tends to make the ligands act as more powerful and potent bacterial agents, thus killing of the more bacteria than the ligand. It is observed that in complexes the positive charge of the metal partially shared with the donor atoms present in the ligand and there may be ␲-electron delocalization over the hole chelate ring. The antifungal screening of the ligand and complexes was performed in vitro against Aspergilus niger, Aspergilus alternaria and Synchytrium endobioticum by the poisonous food technique [23,24]. Potato dextrose agar was used as medium for test. The compounds under study were dissolved in DMF to prepare 0.1% solution of each compound. The solution of compound was mixed with medium and the mixture was poured on a glass petri-dish. When it solidified the plated were inoculated. The inoculated plates were incubated for 48–72 h at 40 ◦ C (±2 ◦ C). The percentage inhibition of mycelial growth of the text fungus was calculated by using the formula: %Inhibition = (C − T ) × 100/C, where C is the diameter of the fungal colonies in the control, T is the diameter of the fungal colonies in treated plates. Results shown in Graph 1, indicates that cobalt(II) complexes are more active as compared to Ni(II) and Cu(II) complexes to inhibit the fungal growth.

Acknowledgement We are thankful to the Principal, Zakir Husain College for providing laboratory facilities and the University Grant Commission, New Delhi for financial assistance RSIC-IIT, Bombay for recording EPR spectra and Dr. Mordhwaj ACBR, University of Delhi, for recording infrared spectra.

[1] K.R. Adam, M. Antolovich, D.S. Baldwin, P.A. Duckworth, A.J. Leong, L.F. Lindoy, M. McPartlin, P.A. Tasker, J. Chem. Soc., Dalton Trans. (1993) 1013. [2] S. Chandra, L.K. Gupta, Spectrochim. Acta 60A (2004) 1563; O.D. Fox, M.G.B. Drew, E.J.S. Wilkinson, P.D. Beer, Chem. Commun. (2000) 319. [3] E. Labisbal, A. Sousa, A. Castineiras, A. Gracia-Vazquez, J. Romero, D.X. West, Polyhedron 19 (2000) 1255. [4] S. Chandra, L.K. Gupta, Spectrochim. Acta 60A (2004) 1751; M. Canadas, E. Lopez-Torres, A.M. Arias, M.A. Mendiola, M.T. Sevilla, Polyhedron 19 (2000) 2059. [5] D.X. West, E. Liberta, S.B. Padhye, R.C. Chikate, P.B. Sonawane, A.S. Kumbar, R.S. Yeranda, Coord. Chem. Rev. 123 (1993) 49. [6] S. Chandra, L.K. Gupta, D. Jain, Spectrochim. Acta 60A (2004) 2411; C.M. Liu, G. Xiong, X.Z. You, Y.J. Liu, Polyhedron 15 (1996) 4565. [7] J.A. Streeky, D.G. Pillsburg, D.H. Busch, Inorg. Chem. 19 (1980) 3148. [8] A.A. Isse, A. Gennaro, E. Vianello, J. Chem. Soc., Dalton Trans. (1993) 2091. [9] S. Chandra, L.K. Gupta, Spectrochim. Acta 60A (2004) 2767; P.D. Beer, N. Berry, M.G.B. Drew, O.D. fox, M.E. Padilla-Tosta, S. Patell, Chem. Commun. (2001) 199. [10] S. Chandra, L.K. Gupta, Spectrochim. Acta 60A (2004) 3079; A. Arquero, M.A. Mendiola, P. Souza, M.T. Sevilla, Polyhedron 15 (1996) 1657. [11] K. Nakamoto, Infrared and Raman Spectra of Coordination Compounds, Wiley Interscience, New York, 1970. [12] S. Chandra, Sangeetika, Aruna Rathi, J. Saudi Chem. Soc. 5 (2001) 175. [13] S. Chandra, Sangeetika, Aruna Rathi, J. Saudi Chem. Soc. 6 (2001) 237. [14] A.B.P. Lever, Inorganic Electronic Spectroscopy, second ed., Elsevier, Amsterdam, 1984. [15] S. Chandra, S.D. Sharma, J. Indian Chem. Soc. 79 (2002) 495. [16] S. Chandra, Sangeetika, Ind. J. Chem. 41A (2002) 1629. [17] B.J. Hathaway, in: J.N. Bradley, R.D. Gillard (Eds.), Essays in Chemistry, Academic Press, New York, 1971. [18] S. Chandra, N. Gupta, L.K. Gupta, Synth. React. Inorg. Met. Org. Chem. 34 (2004) 919; A.A. Isse, A. Gennaro, E. Vianello, J. Chem. Soc., Dalton Trans (1993). [19] E. Pereira, L. Gomez, B. Castro, Inorg. Chim. Acta 271 (1998) 83; S. Chandra, L.K. Gupta, J. Saudi Chem. Soc. 8 (2004) 85. [20] S. Chandra, L.K. Gupta, Sangeetika, Synth. React. Inorg. Met. Org. Chem. 34 (2004) 1591; A. Arquero, M.A. Mendiola, P. Souza, M.T. Sevilla, Polyhedron 15 (1996) 1657. [21] S. Chandra, L.K. Gupta, D. Shukla, J. Saudi Chem. Soc. 8 (2003) 331; P. Bindu, M.R.P. Kurup, T.R. Satyakeerty, Polyhedron 18 (1998) 321. [22] N. Raman, A. Kulandaisamy, A. Shunmugasundaram, K. Jeyasubramanian, Trans. Met. Chem. 26 (2001) 131. [23] A. Bansal, R.V. Singh, J. Ind. Chem. Soc. 40A (2001) 989. [24] W.G. Hanna, M.M. Maowad, Trans. Met. Chem. 26 (2001) 644. [25] E. Berejo, R. Carballo, A. Castineiras, R. Dominguez, C. MaichleMossmer, J. Strahle, D.X. Best, Polyhedron 18 (1999) 3695.