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Nickel(II) complexes with different chromospheres containing macrocyclic ligands: Spectroscopic and electrochemical studies
Spectrochimica Acta Part A 65 (2006) 215–220
Nickel(II) complexes with different chromospheres containing macrocyclic ligands: Spectroscopic and electrochemical studies Sulekh Chandra a , Rajiv Kumar a,b,∗ , Rajeev Singh b a
Department of Chemistry, Zakir Husain College, JL Nehru Marg, University of Delhi, New Delhi 110002, India b Department of Chemistry, University of Delhi, Delhi 110007, India Received 9 September 2005; received in revised form 9 September 2005; accepted 15 October 2005
Abstract The mixed donor tetradentate (L1 = N2 O2 ) and pentadentate (L2 = N2 O2 S) ligands have been prepared by the interaction of 1,3-diaminopropane and thiodiglycolic acid with diamine. These ligands possess two dissimilar coordination sites. Different types of complexes were obtained which have different stoichiometry depending upon the type of ligands. Their structural investigation have been based on elemental analysis, magnetic moment and spectral (ultraviolet, infrared, 1 H NMR, 13 C NMR and mass spectroscopy methods). The Ni(II) complexes show magnetic moments corresponding to two unpaired electrons except [Ni(L1 )](NO3 )2 which is diamagnetic. Ligand field parameters of these complexes were compared. N2 O2 S donor ligand complexes show higher values of ligand field parameters, which are used to detect their geometries. The redox properties and stability of the complexes toward oxidation waves explored by cyclic voltammetry are related to the electronwithdrawing or releasing ability of the substituents of macrocyclic ligands moiety. The Ni(II) complexes displayed Ni(II)/Ni(I) couples irreversible waves associated with Ni(III)/Ni(II) process. © 2005 Elsevier B.V. All rights reserved. Keywords: Macrocyclic ligands; EPR; IR; Cyclic voltammetry; Metal complexes
1. Introduction A major portion of inorganic chemistry may be related to its main branch that is coordination chemistry, which is the investigation of the properties, structures and reactions of complexes formed by ligands coordinated to a transition metal center. Transition metal complexes play a central role in the construction of molecular materials, which display unusual conducting and magnetic properties and find applicability in material chemistry, supramolecular and biochemistry [1–4]. Nickel(II) complexes of macrocyclic ligands are well known to be biologically important and interesting because of their anticarcinogenic, antibacterial and antifungal properties [5]. Also, they have been screened for their medicinal properties because they posses some degree of cytotoxic activity [6]. The electron transfer events of metal ions have enormous importance in metalloenzyme systems, especially respiratory enzymes. The redox properties include oxidation and reductions ∗
Corresponding author. Tel.: +91 1234276530; fax: +91 1234276530. E-mail address: chemistry [email protected] (R. Kumar).
1386-1425/$ – see front matter © 2005 Elsevier B.V. All rights reserved. doi:10.1016/j.saa.2005.10.033
of the central metal ion with ligands have been previously studies and reported [7]. The redox potential of the Ni(I)/Ni(II) and Ni(II)/Ni(III) couples have been shown to be markedly affected by the nature of the solvent, background electrolyte and the structure of the chelating ligand with the complexes [8]. In this article we reported the coordination behavior and characterization of macrocyclic ligands with their nickel(II) metal ion. We have taken in to account the factors which affect the structure of complexes and their geometries and determining the structures based around the theory of magnetic moments and other spectral studies. In this present paper we report the synthesis and characterization of two series of nickel(II) complexes with these macrocyclic ligands (Fig. 1) viz. L1 is 1,7-diaza-10,14-dioxa-4-thia8,9:15,16-dibenzocyclohexadeca-2,6-dione and L2 is l,5-diaza8,12-dioxa-6,7:13,14-dibenzo-cyclotetradecane. 2. Experimental All the chemicals used in the present investigation were of AR grade, purchased from Sigma–Aldrich Co., USA.
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3. Synthesis of macrocyclic ligands 3.1. Preparation of diamines This diamine was prepared by the method reported earlier [2]. 3.2. Preparation of macrocyclic ligands A hot (∼50 ◦ C) aqueous ethanolic solution (20 mL) with calculated amount of corresponding diamines (0.1 mmol) and an ethanolic solution (20 mL) of an equimolar amount of corresponding thiodiglycolic acid or dichloropropane (0.1 mmol) were mixed in the molar ratio 1/1 in 100 mL round bottle flask. The reactant mixture was refluxed for several hours at a temperature of ∼80 ◦ C and the progress of the reaction was ascertained by noting the liberated water azeotropically. The excess of the solvent was first distilled off and then removed using a vacuum pump. On cooling the contents to a temperature at ∼5 ◦ C, the p.p.t. was separated. This was filtered, washed with 50% ethanol and dried over P4 O10 in vacuum desiccators. The proposed chemical structures of the prepared macrocyclic ligands have a good agreement with the stoichiometries concluded from their analytical data and mass spectra and confirmed from the IR spectral data. One of the both ligand was white (L1 ) and other (L2 ) was pale yellow and have melting point below 184 ◦ C. In the IR spectra of these ligands characteristic bands corresponding to >NH, >C O and CH2 O CH2 groups were observed which indicate that these ligands are cyclic in nature.
The mass spectrum of L1 shows a peak at 297 corresponding to the molecular ion (M+ + 1). EIMS m/z (%) 297 (M+ , 72%). The mass spectrum of L2 shows a peak at 371 corresponding to the molecular ion (M+ + 1). EIMS m/z (%) 371 (M+ , 61%). 1 H NMR: (CDCl ) δ: 7.0 (2H, d), 7.1 (2H, m), 6.5 3 (4H, m), 6.80 (2H, d, J = 7.6), 3.7 (4H, m, NH CH2 ), 3.6 (4H, m, O CH2 ), 3.2 (4H, s), 3.4 (4H, CH2 ) and 13 C NMR (CDCl ): 121.01–122.04, 122.10–124.50 (C H ); 3 6 4 144.50–146.0 to 148.0–150.5 (C6 H4 ); 150–152 (NH C6 H4 ); 61.2 ( CH2 ); 62.5–66.20 ( CH2 CH2 CH2 ) for L1 (Fig. 2). 1 H NMR: (CDCl ) δ: 7.2 (4H, m), 6.8 (2H, d, J = 7.1), 6.5 3 (2H, d, J = 6.7), 3.8 (2H, m, NH); 3.0 (4H, m), 2.8 (6H, s) (Fig. 2). 13 C NMR (CDCl ): 122.5–125.5 (C H ); 136.5–137.0 3 6 4 (C6 H4 ); 152.0–152.5 (NH C6 H4 ); 122.0–124.5 (C6 H4 ); 170.5–172.1 (C O); 30.5 ( S CH2 ) for L2 . 3.3. Preparation of Ni(II) complexes A general method has been adopted for the preparation of the complexes. All complexes were prepared by adding stoichiometric quantities of an equimolar amount of corresponding macrocyclic ligand (1 mmol) in a hot (∼45 ◦ C) aqueous ethanol solution (20 mL) and corresponding metal salts (1 mmol) in bi-distilled water and were mixed in the 100 mL round bottle flask. The mixture was refluxed for ∼6 h at a temperature of ∼75 ◦ C on a water bath. The excess of the solvent was first distilled off and then removed using a vacuum pump. On cooling the contents to a temperature of ∼5 ◦ C, the complexes separated. The precipitates formed were filtered off, washed with hot ethanol, and finally dried over P4 O10 in vacuum desiccators. 4. Physical measurements Elemental analysis (CHN) of these complexes was carried out on a Carlo-Erba 1106 elemental analyzer. Molar conductance is measured on an ELICO conductivity bridge (Type CM82T). Magnetic susceptibility measurements were made on Gouy Balance at room temperature using CuSO4 ·5H2 O as calibrant. Molecular weight of the complexes was determined in benzene (freezing point). IR spectra were recorded on a PerkinElmer 137 instrument as nujol mulls/KBr pellets. 1 H NMR
Fig. 1. Suggested structures of ligand L1 and L2 .
Fig. 2.
13 C
NMR spectrum of macrocyclic ligand L1 .
S. Chandra et al. / Spectrochimica Acta Part A 65 (2006) 215–220
217
and 13 NMR spectra of the macrocyclic ligands were recorded on Bruker Avance 300 spectrophotometer at 100 kHz modulation and higher frequency. Electronic spectra were recorded in DMSO solution on a Shimadzu UV mini-1240 spectrophotometer. Electron impact mass spectra were recorded on JEOL, JMS, DX-303 mass spectrometer. Tetramethyl ammonium perchlorate (TEAP) was applied as electrode (AgNO3 in 0.1 M TEAP in DMF solution resulting Ag/Ag+ electrode) on a X–Y HoustonOmmigraphic 2000 recorder Ag/Ag+ . 5. Result and discussion 5.1. IR spectra of the ligands with their nickel(II) complexes
Fig. 3. IR spectrum of [Ni(L1 ) (NCS)2 ].
The most important IR bands of the both ligands and their nickel(II) complexes are summarized in Table 1. The modes of coordination were detected on the basis of following evidences. The >NH band vibrations shifted towards lower frequency on complexation. The phenolic oxygen of the ligands is also showed shifting by 10–15 cm−1 towards the lower frequency region after complexation [9]. Some new characteristic bands at 345–385, 410–430 and 505–530 cm−1 are assigned to ν(Ni–S), ν(Ni–N) and ν(Ni–O) were also observed which confirmed their coordination. 5.2. Bands due to anions The IR spectrum of [Ni(L1 )NCS)2 ] shows ν(CN) band at 2089 cm−1 corresponding to coordinated NCS but [Ni(L2 )(NCS)] showed two medium to strong bands, one at 2180 cm−1 corresponding to coordinate NCS and second at 2015 cm−1 which corresponding to uncoordinate nature of NCS [11]. The infrared spectra of the nitrato complex of L1 and L2 showed bands in range 1385–1386 cm−1 corresponds to an uncoordinated nitrate group (Figs. 3–5).
Fig. 4. IR spectrum of [Ni(L2 )](NCS)(NCS).
5.4. Magnetic moments The magnetic moment for a transition metal compound can generally be directly related to the number of unpaired electrons for the metal ion. This is particularly true for first row transition metals where the experimentally determined magnetic
5.3. Nickel(II) complexes composition The isolated complexes are completely soluble in DMSO. The elemental analysis and molecular formula are listed in Table 3. All of the nickel(II) complexes have the general composition [Ni(L1 )](NO3 )2 , [Ni(L1 )(NCS)2 ] and [Ni(L2 )(NCS)](NCS), [Ni(L2 )](NO3 )2 . Table 1 Important IR absorption bands (cm−1 ) of the complexes with their ligands Complex
νNH
νC–S–C
νC–O–C
νNi–N
νNi–O
νNi–S
C18 H22 N2 O2 [Ni(L1 )](NO3 )2 [Ni(L1 )(NCS)2 ] C19 H20 N2 O4 S [Ni(L2 )](NO3 )2 [Ni(L1 )(NCS)](NCS)
3270 3265 3260 3295 3290 3289
665 657 660 662 652 648
1255 1245 1240 1258 1249 1242
– 423 410 – 421 422
524 530 528 529 526 512
– – – 370 345 385
Fig. 5. IR spectrum of [Ni(L1 )](NO3 )2 .
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Table 2 Magnetic moment and electronic spectral data of the nickel(II) complexes with possible transitions Complexes
ν1
ν2
ν3
Magnetic moments
Transition respectively
NiL1 (NO
18518 13333 14085 14084
19607 18182 16949 18181
27717 34483 45455 35508
Diamagnetic 2.90 3.00 2.92
1A
3 )2 NiL1 (NCS)2 NiL2 (NO3 )2 NiL2 (NCS)2
moment was usually very similar to the spin-only value, i.e., µs.o. = 2[S(S + 1)]1/2 . The magnetic moments of the complexes were recorded at room temperature (300 K) and found in the range of 2.90–3.00 B.M. [12] corresponding to two unpaired electrons, except for [Ni(L1 )](NO3 )2 which was found to be diamagnetic (no unpaired electrons). 5.5. Molar conductance measurements The molar conductance measurements of [Ni(L1 )(NCS)2 ] and [Ni(L1 )](NO3 )2 complexes were 08 and −1 186 cm2 mol−1 corresponding to non-electrolytes and 1:2 electrolyte, respectively [10]. But [Ni(L2 )(NCS)](NCS) and [Ni(L2 )](NO3 )2 complexes showed molar conductance 90 and 185 −1 cm2 mol−1 corresponding to 1:1 and 1:2 electrolytes (Figs. 6 and 7).
Fig. 6. Suggested structures of nickel(II) complexes with L1 .
→ 1 A2g 1 1g → B1g (ν1 ) 1 A → 1 E 3 A → 3 T (F) (ν ) 1g g 2g 2g 2 3 A → 3 T (F) 3 A → 3 T (P) (ν ) Refs. [14–16] 2g 1g 2g 1g 3 1g
1A
5.6. Electronic spectra The electronic spectral values of these complexes are given in Table 2 with their transitions and possible geometries (Figs. 8 and 9). The values spectrochemical parameters Dq of the complexes were found in the range (9.80–10.3) × 103 cm−1 close to that observed for N2 O2 S chromophore containing nickel(II) complexes which had reasonable for most complexes containing oxygen, nitrogen and sulphur. The observed Dq values provided important information about the C O C, C S C and C NH C donor sites of the ligands which coordinates to the nickel(II) ion in different geometries corresponding to the macrocyclic ligand cavities. Thus, it seems reasonable to presume that most of the [Ni(L1 )(NCS)2 ] and [Ni(L1 )(NCS)](NCS) complexes have
Fig. 7. Suggested structures of nickel(II) complexes with L2 .
S. Chandra et al. / Spectrochimica Acta Part A 65 (2006) 215–220
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7. Cyclic voltammetry
Fig. 8. UV spectrum of [Ni(L1 )](NO3 )2 .
N2 O2 or N2 O2 S chromophores. These values of the complexes were equal to the previously reported for the d8 –octahedral complexes and higher than diamagnetic squareplanar [Ni(L1 )](NO3 )2 complex [13]. Suggested structures of the complexes are given in Fig. 4. 6. Ligand field parameters xy
Hamiltonian parameters Ds (500–841), Dt (310–521), Dq (1252–1524) and Dqz (600–710) are the splitting parameters for in-plane and for axial ligands, B values of the studied complexes in the range 69.9–72.0% that of the free ion (Bo = 1030 cm−1 ). The nephelauxetic parameter is in the range 0.66–0.74, indicating that these ligands are in the middle of the nephelauxetic of other nitrogen, sulphur and oxygen donor series. The values are higher than that observed for N6 or O6 chromophore containing nickel(II) complexes confirming the coordination via N, O and S [16,17]. McClure’s parameters ␦t2g and ␦eg depend upon the splitting of the t2g and eg orbitals. ␦ (210–410) and ␦ (315–452) were determined from the crystal field parameters from relation. It indicates that the -bonding effect is relatively more important along the axial direction than -bonding.
Fig. 9. UV spectra of A for [Ni(L1 )(NCS)2 ], B for [Ni(L2 )(NCS)](NCS) and C for [Ni(L2 )](NO3 )2 .
The effect of ring size and donor ability of the ligands with copper(II) complexes have been reported earlier [2]. The binding effect of nickel(II) with macrocyclic ligands cavities and environment show two well-defined electrode couples. The cyclic voltammetry behavior of showing the complexes are found similar displayed two well-defined electrode couples. The complexes showed two successive one electron processes. The first electrode couple is assigned to the irreversible couple Ni(II)/Ni(I) with E1/2 of −1.48 to −1.60 V and represented as follows Ni(II)L + e− → Ni(I)L. The second electrode couple with (E1/2 ) 0.40–0.48 V is assigned Ni(II)/Ni(III) which is a for the irreversible process and have a couple to the peak to peak potential separation (E) with this electrode couple. Ni(II)/Ni(III) is observed with increasing scan rate [4,18]. Thus, the electron transfer process is irreversible and the species formed initially in the electrode process may also react further to give products that are not deoxidized at the same potential as in the first formed species [19,20]. Ep values indicate that the potential differences between original and reduced metal species. On comparing the reduction behaviors of L1 and L2 complexes, it was concluded that the electrical property of the metal is affected directly by ring size and unsaturated behavior of ligands cavity. It is also observed that oxidation potential of metal is also slightly affected by anions. It was also observed that the macrocyclic ligands moiety has a significant effect on E1/2 for all the complexes; electronwithdrawing groups stabilize the nickel(II) in the complexes while the electron-donating group favor oxidation to nickel(III). This is possibly because the electron-withdrawing anion makes the complex more positively charged and it favor the reduction of metal ion. Similarly the electron-donating groups make the complexes less positively charged. Electro potentials of Ni(II)/Ni(III) couple is showing sensitivity to the nature of donor atoms in macrocyclic ligands moiety. 8. Conclusion A series of nickel(II) complexes with macrocyclic ligands have been prepared and fully characterized. The characteristics of the prepared complexes are discussed. The coordination behavior of the anions is also discussed on the basis of IR and molar conductance measurements. It was observed that coordination of anions with metal ions was effected by number of coordination sites, which was further confirmed by molar conductance. The cyclic voltammetric of the complexes showed that the nickel(II) compounds undergo one-electron reduction and oxidation to form the corresponding to nickel(I) and nickel(III) species. These kinds of complexes may be found being used as one-electron redox reagents since the former is a strong reducing agent and the latter is strong oxidizing agent. The geometries of the complexes are also effected by number of coordination sites and metal oxidation states. Ligand L1 has four coordination sites and formed different geometrical complexes with different anions. The important characteristics of
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Table 3 Elemental analysis data found (calculated)
References
Complex
[1] E.I. Solomon, T.C. Brunold, M.Z. Davis, J.N. Kemseley, S.K. Lee, N. Lehnert, A.J. Skulan, Y.S. Yang, J. Zhou, Chem. Rev. 100 (2000) 235. [2] S. Chandra, R. Kumar, Trans. Met. Chem. 29 (3) (2003) 269. [3] S. Chandra, R. Kumar, Spectrochim. Acta Part A 61 (2005) 437. [4] S. Chandra, R. Kumar, Trans. Synth. React. Inorg. Met-Org. Nano-Met. Chem. 35 (2005) 103. [5] S. Chandra, R. Kumar, Trans. Synth. React. Inorg. Met-Org. Nano-Met. Chem. 35 (2005) 161. [6] N.P. Farrell, Uses of Inorganic Chemistry in Medicine, Royal Society of Chemistry, Cambridge, UK, 1999. [7] M.S. El-Shahawi, W.E. Smith, Analyst 119 (1994) 327. [8] I. Schmidt, P.J. Chmielewski, Inorg. Chem. 42 (18) (2003) 5579. [9] R.M. Silverstein, G.C. Bassler, T.C. Movril, Spectroscopic Identification of Organic Compounds, fourth ed., Wiley, New York, 1981. [10] M. Tosha, R.G. Barclay, G.H. Robin, T.M. Lemaire, K.T. Laurence, Chem. Commun. 2141 (2000). [11] K. Nakamoto, Infrared and Raman Spectra of Inorganic and Coordination Compounds, vol. 305, third ed., Wiley, New York, 1978. [12] S.V. Kryatov, B.S. Mohanraj, V.V. Tarasov, O.P. Kryatova, E.V. RybakAkimova, B. Nuthakki, J.F. Rusling, R.J. Staples, A.Y. Nazarenko, Inorg. Chem. 41 (4) (2002) 923. [13] M.P. Shores, J.J. Sokol, J.R. Long, Am. Chem. Soc. 124 (10) (2002) 2279. [14] A.I. Vogel, A Text Book of Quantitative Inorganic Analysis, Longmans, London, 1961. [15] A.B.P. Lever, Inorganic Electronic Spectroscopy, Elsevier, Amsterdam, 1986. [16] R. Kr¨amer, L. Kovbasyuk, H. Pritzko, New J. Chem. 5 (2002) 516. [17] C.-Y. Hua, W. Li, L.-Y. Dai, Y.-K. Shan, J. Chem. Res. 2 (2004) 103. [18] I. Zilbermann, J. Hayon, T. Katchalski, O. Raveh, J. Rishpon, A.I. Shames, A. Bettelheim, J. Electro. Chem. Soc. 144 (1999) 228. [19] S.-G. Kang, K.K. Kweon, S.K. Jung, Bull. Korean Chem. Soc. 12 (1991) 483. [20] A.J. Bard, L.R. Faulkner, Electrochemical Methods Fundamental and Applications, John Wiley, New York, 1980.
C18 H22 N2 O2 [Ni(L1 )](NO3 )2 NiC18 H22 N4 O8 [Ni(L1 )(NCS)2 ] NiC20 H22 N4 O2 S2 C19 H20 N2 O4 S [Ni(L2 )](NO3 )2 NiC19 H20 N4 O10 S [Ni(L1 )(NCS)](NCS) NiC21 H20 N4 O4 S3
Elemental analysis calculated (found) %C
%H
%N
%Ni
72.30 (72.46) 44.90 (44.94) 50.70 (50.76) 61.20 (61.27) 41.00 (41.11) 46.00 (46.09)
7.38 (7.43) 4.54 (4.61) 4.61 (4.69) 5.06 (5.41) 3.51 (3.63) 3.52 (3.60)
9.32 (9.39) 11.51 (11.65) 11.70 (11.84) 7.39 (7.52) 11.00 (11.09) 10.15 (10.24)
– 12.10 (12.20) 12.36 (12.40) – 10.47 (10.57) 10.52 (10.72)
nitrate ions with both ligands are different form the NCS complexes. NO3 complexes of both ligands are compared and it is observe that the complex with L1 is square plannar but with L2 , it is a five coordinate arrangement in macrocyclic cavities. NCS complexes of both ligands are compared and it observes that these complexes are six coordinated. It proved direct effect of coordination sites on complexeation (Table 3). Acknowledgements One of the authors (Rajiv Kumar) gratefully acknowledges his younger brother Bitto for motivation. Thanks to the University Grants Commission (UGC), New Delhi for financial assistance and the University Science Instrumentation Center, Delhi University, for recording IR spectra. Thanks are also due to the Solid State Physics Laboratory India for recording magnetic moments.
Nickel(II) complexes with different chromospheres containing macrocyclic ligands: Spectroscopic and electrochemical studies Sulekh Chandra a , Rajiv Kumar a,b,∗ , Rajeev Singh b a
Department of Chemistry, Zakir Husain College, JL Nehru Marg, University of Delhi, New Delhi 110002, India b Department of Chemistry, University of Delhi, Delhi 110007, India Received 9 September 2005; received in revised form 9 September 2005; accepted 15 October 2005
Abstract The mixed donor tetradentate (L1 = N2 O2 ) and pentadentate (L2 = N2 O2 S) ligands have been prepared by the interaction of 1,3-diaminopropane and thiodiglycolic acid with diamine. These ligands possess two dissimilar coordination sites. Different types of complexes were obtained which have different stoichiometry depending upon the type of ligands. Their structural investigation have been based on elemental analysis, magnetic moment and spectral (ultraviolet, infrared, 1 H NMR, 13 C NMR and mass spectroscopy methods). The Ni(II) complexes show magnetic moments corresponding to two unpaired electrons except [Ni(L1 )](NO3 )2 which is diamagnetic. Ligand field parameters of these complexes were compared. N2 O2 S donor ligand complexes show higher values of ligand field parameters, which are used to detect their geometries. The redox properties and stability of the complexes toward oxidation waves explored by cyclic voltammetry are related to the electronwithdrawing or releasing ability of the substituents of macrocyclic ligands moiety. The Ni(II) complexes displayed Ni(II)/Ni(I) couples irreversible waves associated with Ni(III)/Ni(II) process. © 2005 Elsevier B.V. All rights reserved. Keywords: Macrocyclic ligands; EPR; IR; Cyclic voltammetry; Metal complexes
1. Introduction A major portion of inorganic chemistry may be related to its main branch that is coordination chemistry, which is the investigation of the properties, structures and reactions of complexes formed by ligands coordinated to a transition metal center. Transition metal complexes play a central role in the construction of molecular materials, which display unusual conducting and magnetic properties and find applicability in material chemistry, supramolecular and biochemistry [1–4]. Nickel(II) complexes of macrocyclic ligands are well known to be biologically important and interesting because of their anticarcinogenic, antibacterial and antifungal properties [5]. Also, they have been screened for their medicinal properties because they posses some degree of cytotoxic activity [6]. The electron transfer events of metal ions have enormous importance in metalloenzyme systems, especially respiratory enzymes. The redox properties include oxidation and reductions ∗
Corresponding author. Tel.: +91 1234276530; fax: +91 1234276530. E-mail address: chemistry [email protected] (R. Kumar).
1386-1425/$ – see front matter © 2005 Elsevier B.V. All rights reserved. doi:10.1016/j.saa.2005.10.033
of the central metal ion with ligands have been previously studies and reported [7]. The redox potential of the Ni(I)/Ni(II) and Ni(II)/Ni(III) couples have been shown to be markedly affected by the nature of the solvent, background electrolyte and the structure of the chelating ligand with the complexes [8]. In this article we reported the coordination behavior and characterization of macrocyclic ligands with their nickel(II) metal ion. We have taken in to account the factors which affect the structure of complexes and their geometries and determining the structures based around the theory of magnetic moments and other spectral studies. In this present paper we report the synthesis and characterization of two series of nickel(II) complexes with these macrocyclic ligands (Fig. 1) viz. L1 is 1,7-diaza-10,14-dioxa-4-thia8,9:15,16-dibenzocyclohexadeca-2,6-dione and L2 is l,5-diaza8,12-dioxa-6,7:13,14-dibenzo-cyclotetradecane. 2. Experimental All the chemicals used in the present investigation were of AR grade, purchased from Sigma–Aldrich Co., USA.
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3. Synthesis of macrocyclic ligands 3.1. Preparation of diamines This diamine was prepared by the method reported earlier [2]. 3.2. Preparation of macrocyclic ligands A hot (∼50 ◦ C) aqueous ethanolic solution (20 mL) with calculated amount of corresponding diamines (0.1 mmol) and an ethanolic solution (20 mL) of an equimolar amount of corresponding thiodiglycolic acid or dichloropropane (0.1 mmol) were mixed in the molar ratio 1/1 in 100 mL round bottle flask. The reactant mixture was refluxed for several hours at a temperature of ∼80 ◦ C and the progress of the reaction was ascertained by noting the liberated water azeotropically. The excess of the solvent was first distilled off and then removed using a vacuum pump. On cooling the contents to a temperature at ∼5 ◦ C, the p.p.t. was separated. This was filtered, washed with 50% ethanol and dried over P4 O10 in vacuum desiccators. The proposed chemical structures of the prepared macrocyclic ligands have a good agreement with the stoichiometries concluded from their analytical data and mass spectra and confirmed from the IR spectral data. One of the both ligand was white (L1 ) and other (L2 ) was pale yellow and have melting point below 184 ◦ C. In the IR spectra of these ligands characteristic bands corresponding to >NH, >C O and CH2 O CH2 groups were observed which indicate that these ligands are cyclic in nature.
The mass spectrum of L1 shows a peak at 297 corresponding to the molecular ion (M+ + 1). EIMS m/z (%) 297 (M+ , 72%). The mass spectrum of L2 shows a peak at 371 corresponding to the molecular ion (M+ + 1). EIMS m/z (%) 371 (M+ , 61%). 1 H NMR: (CDCl ) δ: 7.0 (2H, d), 7.1 (2H, m), 6.5 3 (4H, m), 6.80 (2H, d, J = 7.6), 3.7 (4H, m, NH CH2 ), 3.6 (4H, m, O CH2 ), 3.2 (4H, s), 3.4 (4H, CH2 ) and 13 C NMR (CDCl ): 121.01–122.04, 122.10–124.50 (C H ); 3 6 4 144.50–146.0 to 148.0–150.5 (C6 H4 ); 150–152 (NH C6 H4 ); 61.2 ( CH2 ); 62.5–66.20 ( CH2 CH2 CH2 ) for L1 (Fig. 2). 1 H NMR: (CDCl ) δ: 7.2 (4H, m), 6.8 (2H, d, J = 7.1), 6.5 3 (2H, d, J = 6.7), 3.8 (2H, m, NH); 3.0 (4H, m), 2.8 (6H, s) (Fig. 2). 13 C NMR (CDCl ): 122.5–125.5 (C H ); 136.5–137.0 3 6 4 (C6 H4 ); 152.0–152.5 (NH C6 H4 ); 122.0–124.5 (C6 H4 ); 170.5–172.1 (C O); 30.5 ( S CH2 ) for L2 . 3.3. Preparation of Ni(II) complexes A general method has been adopted for the preparation of the complexes. All complexes were prepared by adding stoichiometric quantities of an equimolar amount of corresponding macrocyclic ligand (1 mmol) in a hot (∼45 ◦ C) aqueous ethanol solution (20 mL) and corresponding metal salts (1 mmol) in bi-distilled water and were mixed in the 100 mL round bottle flask. The mixture was refluxed for ∼6 h at a temperature of ∼75 ◦ C on a water bath. The excess of the solvent was first distilled off and then removed using a vacuum pump. On cooling the contents to a temperature of ∼5 ◦ C, the complexes separated. The precipitates formed were filtered off, washed with hot ethanol, and finally dried over P4 O10 in vacuum desiccators. 4. Physical measurements Elemental analysis (CHN) of these complexes was carried out on a Carlo-Erba 1106 elemental analyzer. Molar conductance is measured on an ELICO conductivity bridge (Type CM82T). Magnetic susceptibility measurements were made on Gouy Balance at room temperature using CuSO4 ·5H2 O as calibrant. Molecular weight of the complexes was determined in benzene (freezing point). IR spectra were recorded on a PerkinElmer 137 instrument as nujol mulls/KBr pellets. 1 H NMR
Fig. 1. Suggested structures of ligand L1 and L2 .
Fig. 2.
13 C
NMR spectrum of macrocyclic ligand L1 .
S. Chandra et al. / Spectrochimica Acta Part A 65 (2006) 215–220
217
and 13 NMR spectra of the macrocyclic ligands were recorded on Bruker Avance 300 spectrophotometer at 100 kHz modulation and higher frequency. Electronic spectra were recorded in DMSO solution on a Shimadzu UV mini-1240 spectrophotometer. Electron impact mass spectra were recorded on JEOL, JMS, DX-303 mass spectrometer. Tetramethyl ammonium perchlorate (TEAP) was applied as electrode (AgNO3 in 0.1 M TEAP in DMF solution resulting Ag/Ag+ electrode) on a X–Y HoustonOmmigraphic 2000 recorder Ag/Ag+ . 5. Result and discussion 5.1. IR spectra of the ligands with their nickel(II) complexes
Fig. 3. IR spectrum of [Ni(L1 ) (NCS)2 ].
The most important IR bands of the both ligands and their nickel(II) complexes are summarized in Table 1. The modes of coordination were detected on the basis of following evidences. The >NH band vibrations shifted towards lower frequency on complexation. The phenolic oxygen of the ligands is also showed shifting by 10–15 cm−1 towards the lower frequency region after complexation [9]. Some new characteristic bands at 345–385, 410–430 and 505–530 cm−1 are assigned to ν(Ni–S), ν(Ni–N) and ν(Ni–O) were also observed which confirmed their coordination. 5.2. Bands due to anions The IR spectrum of [Ni(L1 )NCS)2 ] shows ν(CN) band at 2089 cm−1 corresponding to coordinated NCS but [Ni(L2 )(NCS)] showed two medium to strong bands, one at 2180 cm−1 corresponding to coordinate NCS and second at 2015 cm−1 which corresponding to uncoordinate nature of NCS [11]. The infrared spectra of the nitrato complex of L1 and L2 showed bands in range 1385–1386 cm−1 corresponds to an uncoordinated nitrate group (Figs. 3–5).
Fig. 4. IR spectrum of [Ni(L2 )](NCS)(NCS).
5.4. Magnetic moments The magnetic moment for a transition metal compound can generally be directly related to the number of unpaired electrons for the metal ion. This is particularly true for first row transition metals where the experimentally determined magnetic
5.3. Nickel(II) complexes composition The isolated complexes are completely soluble in DMSO. The elemental analysis and molecular formula are listed in Table 3. All of the nickel(II) complexes have the general composition [Ni(L1 )](NO3 )2 , [Ni(L1 )(NCS)2 ] and [Ni(L2 )(NCS)](NCS), [Ni(L2 )](NO3 )2 . Table 1 Important IR absorption bands (cm−1 ) of the complexes with their ligands Complex
νNH
νC–S–C
νC–O–C
νNi–N
νNi–O
νNi–S
C18 H22 N2 O2 [Ni(L1 )](NO3 )2 [Ni(L1 )(NCS)2 ] C19 H20 N2 O4 S [Ni(L2 )](NO3 )2 [Ni(L1 )(NCS)](NCS)
3270 3265 3260 3295 3290 3289
665 657 660 662 652 648
1255 1245 1240 1258 1249 1242
– 423 410 – 421 422
524 530 528 529 526 512
– – – 370 345 385
Fig. 5. IR spectrum of [Ni(L1 )](NO3 )2 .
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Table 2 Magnetic moment and electronic spectral data of the nickel(II) complexes with possible transitions Complexes
ν1
ν2
ν3
Magnetic moments
Transition respectively
NiL1 (NO
18518 13333 14085 14084
19607 18182 16949 18181
27717 34483 45455 35508
Diamagnetic 2.90 3.00 2.92
1A
3 )2 NiL1 (NCS)2 NiL2 (NO3 )2 NiL2 (NCS)2
moment was usually very similar to the spin-only value, i.e., µs.o. = 2[S(S + 1)]1/2 . The magnetic moments of the complexes were recorded at room temperature (300 K) and found in the range of 2.90–3.00 B.M. [12] corresponding to two unpaired electrons, except for [Ni(L1 )](NO3 )2 which was found to be diamagnetic (no unpaired electrons). 5.5. Molar conductance measurements The molar conductance measurements of [Ni(L1 )(NCS)2 ] and [Ni(L1 )](NO3 )2 complexes were 08 and −1 186 cm2 mol−1 corresponding to non-electrolytes and 1:2 electrolyte, respectively [10]. But [Ni(L2 )(NCS)](NCS) and [Ni(L2 )](NO3 )2 complexes showed molar conductance 90 and 185 −1 cm2 mol−1 corresponding to 1:1 and 1:2 electrolytes (Figs. 6 and 7).
Fig. 6. Suggested structures of nickel(II) complexes with L1 .
→ 1 A2g 1 1g → B1g (ν1 ) 1 A → 1 E 3 A → 3 T (F) (ν ) 1g g 2g 2g 2 3 A → 3 T (F) 3 A → 3 T (P) (ν ) Refs. [14–16] 2g 1g 2g 1g 3 1g
1A
5.6. Electronic spectra The electronic spectral values of these complexes are given in Table 2 with their transitions and possible geometries (Figs. 8 and 9). The values spectrochemical parameters Dq of the complexes were found in the range (9.80–10.3) × 103 cm−1 close to that observed for N2 O2 S chromophore containing nickel(II) complexes which had reasonable for most complexes containing oxygen, nitrogen and sulphur. The observed Dq values provided important information about the C O C, C S C and C NH C donor sites of the ligands which coordinates to the nickel(II) ion in different geometries corresponding to the macrocyclic ligand cavities. Thus, it seems reasonable to presume that most of the [Ni(L1 )(NCS)2 ] and [Ni(L1 )(NCS)](NCS) complexes have
Fig. 7. Suggested structures of nickel(II) complexes with L2 .
S. Chandra et al. / Spectrochimica Acta Part A 65 (2006) 215–220
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7. Cyclic voltammetry
Fig. 8. UV spectrum of [Ni(L1 )](NO3 )2 .
N2 O2 or N2 O2 S chromophores. These values of the complexes were equal to the previously reported for the d8 –octahedral complexes and higher than diamagnetic squareplanar [Ni(L1 )](NO3 )2 complex [13]. Suggested structures of the complexes are given in Fig. 4. 6. Ligand field parameters xy
Hamiltonian parameters Ds (500–841), Dt (310–521), Dq (1252–1524) and Dqz (600–710) are the splitting parameters for in-plane and for axial ligands, B values of the studied complexes in the range 69.9–72.0% that of the free ion (Bo = 1030 cm−1 ). The nephelauxetic parameter is in the range 0.66–0.74, indicating that these ligands are in the middle of the nephelauxetic of other nitrogen, sulphur and oxygen donor series. The values are higher than that observed for N6 or O6 chromophore containing nickel(II) complexes confirming the coordination via N, O and S [16,17]. McClure’s parameters ␦t2g and ␦eg depend upon the splitting of the t2g and eg orbitals. ␦ (210–410) and ␦ (315–452) were determined from the crystal field parameters from relation. It indicates that the -bonding effect is relatively more important along the axial direction than -bonding.
Fig. 9. UV spectra of A for [Ni(L1 )(NCS)2 ], B for [Ni(L2 )(NCS)](NCS) and C for [Ni(L2 )](NO3 )2 .
The effect of ring size and donor ability of the ligands with copper(II) complexes have been reported earlier [2]. The binding effect of nickel(II) with macrocyclic ligands cavities and environment show two well-defined electrode couples. The cyclic voltammetry behavior of showing the complexes are found similar displayed two well-defined electrode couples. The complexes showed two successive one electron processes. The first electrode couple is assigned to the irreversible couple Ni(II)/Ni(I) with E1/2 of −1.48 to −1.60 V and represented as follows Ni(II)L + e− → Ni(I)L. The second electrode couple with (E1/2 ) 0.40–0.48 V is assigned Ni(II)/Ni(III) which is a for the irreversible process and have a couple to the peak to peak potential separation (E) with this electrode couple. Ni(II)/Ni(III) is observed with increasing scan rate [4,18]. Thus, the electron transfer process is irreversible and the species formed initially in the electrode process may also react further to give products that are not deoxidized at the same potential as in the first formed species [19,20]. Ep values indicate that the potential differences between original and reduced metal species. On comparing the reduction behaviors of L1 and L2 complexes, it was concluded that the electrical property of the metal is affected directly by ring size and unsaturated behavior of ligands cavity. It is also observed that oxidation potential of metal is also slightly affected by anions. It was also observed that the macrocyclic ligands moiety has a significant effect on E1/2 for all the complexes; electronwithdrawing groups stabilize the nickel(II) in the complexes while the electron-donating group favor oxidation to nickel(III). This is possibly because the electron-withdrawing anion makes the complex more positively charged and it favor the reduction of metal ion. Similarly the electron-donating groups make the complexes less positively charged. Electro potentials of Ni(II)/Ni(III) couple is showing sensitivity to the nature of donor atoms in macrocyclic ligands moiety. 8. Conclusion A series of nickel(II) complexes with macrocyclic ligands have been prepared and fully characterized. The characteristics of the prepared complexes are discussed. The coordination behavior of the anions is also discussed on the basis of IR and molar conductance measurements. It was observed that coordination of anions with metal ions was effected by number of coordination sites, which was further confirmed by molar conductance. The cyclic voltammetric of the complexes showed that the nickel(II) compounds undergo one-electron reduction and oxidation to form the corresponding to nickel(I) and nickel(III) species. These kinds of complexes may be found being used as one-electron redox reagents since the former is a strong reducing agent and the latter is strong oxidizing agent. The geometries of the complexes are also effected by number of coordination sites and metal oxidation states. Ligand L1 has four coordination sites and formed different geometrical complexes with different anions. The important characteristics of
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Table 3 Elemental analysis data found (calculated)
References
Complex
[1] E.I. Solomon, T.C. Brunold, M.Z. Davis, J.N. Kemseley, S.K. Lee, N. Lehnert, A.J. Skulan, Y.S. Yang, J. Zhou, Chem. Rev. 100 (2000) 235. [2] S. Chandra, R. Kumar, Trans. Met. Chem. 29 (3) (2003) 269. [3] S. Chandra, R. Kumar, Spectrochim. Acta Part A 61 (2005) 437. [4] S. Chandra, R. Kumar, Trans. Synth. React. Inorg. Met-Org. Nano-Met. Chem. 35 (2005) 103. [5] S. Chandra, R. Kumar, Trans. Synth. React. Inorg. Met-Org. Nano-Met. Chem. 35 (2005) 161. [6] N.P. Farrell, Uses of Inorganic Chemistry in Medicine, Royal Society of Chemistry, Cambridge, UK, 1999. [7] M.S. El-Shahawi, W.E. Smith, Analyst 119 (1994) 327. [8] I. Schmidt, P.J. Chmielewski, Inorg. Chem. 42 (18) (2003) 5579. [9] R.M. Silverstein, G.C. Bassler, T.C. Movril, Spectroscopic Identification of Organic Compounds, fourth ed., Wiley, New York, 1981. [10] M. Tosha, R.G. Barclay, G.H. Robin, T.M. Lemaire, K.T. Laurence, Chem. Commun. 2141 (2000). [11] K. Nakamoto, Infrared and Raman Spectra of Inorganic and Coordination Compounds, vol. 305, third ed., Wiley, New York, 1978. [12] S.V. Kryatov, B.S. Mohanraj, V.V. Tarasov, O.P. Kryatova, E.V. RybakAkimova, B. Nuthakki, J.F. Rusling, R.J. Staples, A.Y. Nazarenko, Inorg. Chem. 41 (4) (2002) 923. [13] M.P. Shores, J.J. Sokol, J.R. Long, Am. Chem. Soc. 124 (10) (2002) 2279. [14] A.I. Vogel, A Text Book of Quantitative Inorganic Analysis, Longmans, London, 1961. [15] A.B.P. Lever, Inorganic Electronic Spectroscopy, Elsevier, Amsterdam, 1986. [16] R. Kr¨amer, L. Kovbasyuk, H. Pritzko, New J. Chem. 5 (2002) 516. [17] C.-Y. Hua, W. Li, L.-Y. Dai, Y.-K. Shan, J. Chem. Res. 2 (2004) 103. [18] I. Zilbermann, J. Hayon, T. Katchalski, O. Raveh, J. Rishpon, A.I. Shames, A. Bettelheim, J. Electro. Chem. Soc. 144 (1999) 228. [19] S.-G. Kang, K.K. Kweon, S.K. Jung, Bull. Korean Chem. Soc. 12 (1991) 483. [20] A.J. Bard, L.R. Faulkner, Electrochemical Methods Fundamental and Applications, John Wiley, New York, 1980.
C18 H22 N2 O2 [Ni(L1 )](NO3 )2 NiC18 H22 N4 O8 [Ni(L1 )(NCS)2 ] NiC20 H22 N4 O2 S2 C19 H20 N2 O4 S [Ni(L2 )](NO3 )2 NiC19 H20 N4 O10 S [Ni(L1 )(NCS)](NCS) NiC21 H20 N4 O4 S3
Elemental analysis calculated (found) %C
%H
%N
%Ni
72.30 (72.46) 44.90 (44.94) 50.70 (50.76) 61.20 (61.27) 41.00 (41.11) 46.00 (46.09)
7.38 (7.43) 4.54 (4.61) 4.61 (4.69) 5.06 (5.41) 3.51 (3.63) 3.52 (3.60)
9.32 (9.39) 11.51 (11.65) 11.70 (11.84) 7.39 (7.52) 11.00 (11.09) 10.15 (10.24)
– 12.10 (12.20) 12.36 (12.40) – 10.47 (10.57) 10.52 (10.72)
nitrate ions with both ligands are different form the NCS complexes. NO3 complexes of both ligands are compared and it is observe that the complex with L1 is square plannar but with L2 , it is a five coordinate arrangement in macrocyclic cavities. NCS complexes of both ligands are compared and it observes that these complexes are six coordinated. It proved direct effect of coordination sites on complexeation (Table 3). Acknowledgements One of the authors (Rajiv Kumar) gratefully acknowledges his younger brother Bitto for motivation. Thanks to the University Grants Commission (UGC), New Delhi for financial assistance and the University Science Instrumentation Center, Delhi University, for recording IR spectra. Thanks are also due to the Solid State Physics Laboratory India for recording magnetic moments.