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Polyhedron 26 (2007) 4009–4018

Synthesis, electrochemical and in situ spectroelectrochemical studiesof new transition metal complexes with two new Schiff-bases

containing N2O2/N2O4 donor groups

Ahmet Kilic a,*, Esref Tas a, Bedriye Deveci a, Ismail Yilmaz b,*

a Department of Chemistry, Harran University, Sanliurfa 63510, Turkeyb Department of Chemistry, Technical University of Istanbul, 34469 Istanbul, Turkey

Received 5 February 2007; accepted 1 May 2007Available online 18 May 2007

Abstract

New Schiff-base copper and cobalt complexes, [Cu(L1)], [Cu(L2)] and [Co(L1)], [Co(L2)] (where L1 = N-N 0-bis(3,5-di-tert-butylsalicyl-aldimine)-1,4-cyclohexane bis(methylamine) and L2 = N-N 0-bis(3,5-di-tert-butylsalicylaldimine)-1,8-diamino-3,6-dioxaoctane), were syn-thesized and characterized using elemental analysis, IR spectra, UV–Vis spectra, magnetic susceptibility measurements, 1H and 13CNMR spectroscopy, thermal analysis and molar conductance (KM). Their electro-spectrochemical properties were investigated using cyc-lic voltammetric (CV) and thin-layer spectroelectrochemical techniques in a dichloromethane solution (CH2Cl2). The CV of [Cu(L2)]showed a lower oxidation potential than that of [Cu(L1)] under the same experimental conditions. The oxidation wave (II) of[Cu(L2)] was accompanied by an EC process (II 0), which was not observed for [Cu(L1)]. Also, [Cu(L2)] exhibited a reduction process,but [Cu(L1)] did not. These results indicate that the Cu(II) ion in [Cu(L2)] is coordinated by N2O4 donor sites while [Cu(L1)] presentsa square-planar structure with N2O2 donor sites. Both oxidation processes for [Co(L1)] and [Co(L2)] are based on the cobalt center,and they are assigned to Co(II)/Co(III) couples. The spectroelectrochemical results indicate that the oxidized species of [Cu(L2)] is similarto that of [Cu(L1)], the only difference being that the absorption bands of the oxidized species for [Cu(L2)] shift to lower energy comparedwith those of [Cu(L1)] because of their different coordination environment. The geometry of [Cu(L2)] changed into square-planar afterthe complex was totally oxidized and the neutral complex was only recovered following the EC process, as observed from the CV of[Cu(L2)]. For the two cobalt complexes, the bands corresponding to the p! p* transitions disappeared and new bands with smallred shifts and of lower intensity were observed during the oxidation process. These new bands are attributed to the LMCT transitionas observed in the case of the oxidation processes of the cobalt complexes.� 2007 Elsevier Ltd. All rights reserved.

Keywords: Schiff-base; Complexes; Copper; Cobalt; Electrochemistry; Spectroelectrochemistry

1. Introduction

Schiff’s base ligands have been in the chemistry cata-logue for over 150 years [1]. Their instant and enduringpopularity undoubtedly stem from the ease with whichthey can be synthesized, their bewildering versatility and

0277-5387/$ - see front matter � 2007 Elsevier Ltd. All rights reserved.

doi:10.1016/j.poly.2007.05.013

* Corresponding authors. Tel.: +90 4143440020/1272; fax: +904143440051 (A. Kilic), tel.: +90 2122856831; fax: +90 2122856386(I. Yilmaz).

E-mail addresses: kilica63@harran.edu.tr (A. Kilic), iyilmaz@itu.edu.tr (I. Yilmaz).

their wide ranging complexing ability once formed. The lit-erature clearly shows that the study of this diverse ligandsystem is linked with many of the key advances made ininorganic chemistry. Not only have they played a seminalrole in the development of modern coordination chemistry[2], but they can also be found at key points in the devel-opment of inorganic biochemistry [3], catalysis [4,5] andoptical materials [6]. Although the magnetic, spectroscopicand catalytic properties of these Schiff-base complexes arewell documented [7], it still seems there could be new andspecific applications for such a unique class of compound.

4010 A. Kilic et al. / Polyhedron 26 (2007) 4009–4018

A considerable number of Schiff-base complexes havepotential biological interest, being used as more or lesssuccessful models of biological compounds [8]. Transitionmetal complexes of Schiff-bases containing tetradentateligands have also showed antimicrobial activity [9]. It hasbeen reported that transition metal complexes with redoxactive ligands bearing sterically hindered salicylaldiminesundergo one or two electron transfers [10]. Redox-activetransition metal complexes that stabilize various oxidationstates of the metal centers are of considerable interest,because of their potential significance as models of redoxmetalloenzymes [11,12] and as effective redox reactantsor catalysts [13]. Furthermore, a recently emerged fieldof research concerns the use of redox-active metal com-plexes, for instance, copper(II) complexes with salen andrelated Schiff-bases, as synthetic chemical nucleases orDNA damaging agents [14]. For a specific metal ion, themain factors that determine the redox properties includethe nature and the arrangement of the donor atomsaround the metal bonding site and the nature and the posi-tion of the substituents. Among the various ligands, manySchiff-bases derived from salicylaldehyde and related aro-matic aldehydes have been found to stabilize copper(I),with respect to the Cu(II/I) potential [15]. Although theredox behavior of a number of metal complexes containingSchiff-base ligands have been reported, the electrochemicalproperties of such complexes are not completely clear[16–18]. Several researches have proposed that the redoxpotential of macrocyclic and Schiff-base complexes isdirectly related to many of the biologically relevant chem-ical characteristics of the entire complex, e.g. dioxygenbinding ability and nucleophilicity [19]. Thus, there hasbeen a strong interest in determining thermodynamicallymeaningful redox potentials of copper Schiff-base com-plexes and in understanding the relationship between thesepotentials and the detailed structure of the Schiff-baseligand [20].

In the present paper we report the synthesis and charac-terization of two new Schiff- base ligands (L1 and L2, whereL1 = N-N 0-bis(3,5-di-tert-butylsalicylaldimine)-1,4-cyclo-hexane bis(methylamine) and L2 = N-N 0-bis(3,5-di-tert-butylsalicylaldimine)-1,8-diamino-3,6-dioxaoctane) involv-ing N2O2 and N2O4 donor sites, and their copper(II) andcobalt(II) complexes. It is known that the combination ofelectrochemical and spectroelectrochemical methods pro-vides a power tool to reveal the complementary nature ofthe molecular structure and electrochemical informationof the electroactive molecules [21–26]. Very little workhas been done on spectroelectrochemical characterizationof Schiff-base metal complexes [27,28]. The aim of thisstudy is to establish a comparative electro-spectrochemicalstudy on the new Schiff-base copper and cobalt complexesbased on the different molecular structures with N2O2 andN2O4 donor sites. In this study, the different binding modesof the new Schiff-base ligands for complexing with sometransition metals such as copper and cobalt were confirmedby electrochemical and thin-layer spectroelectrochemical

techniques, in addition to other spectroscopic methodssuch as IR, magnetic susceptibility, 1H and 13C NMR.

2. Experimental

2.1. Material and methods

Unless otherwise stated, all chemicals were of analyticalreagent-grade and were purchased from Sigma, Merck andAcross Organics. 3,5-Di-tert-butyl-2-hidroxybenzaldehyde(3,5-DTB) was synthesized according to the literatureprocedure [29a]. N-N 0-Bis(3,5-di-tert-butylsalicylaldi-mine)-1,4-cyclohexane bis(methylamine) (L1) and N-N 0-bis(3,5-di-tert-butylsalicylaldimine)-1,8-diamino-3,6-dioxa-octane (L2) were prepared here for the first time. The ele-mental analyses and 1H NMR and 13C NMR spectra werecarried out in the laboratory of Tubitak (Scientific andTechnical Research Council of Turkey), IR spectra wererecorded on a Perkin Elmer Spectrum RXI FT-IR Spec-trometer as KBr pellets, Magnetic Susceptibilities weredetermined on a Sherwood Scientific Magnetic Suscepti-bility Balance (Model MK1) at room temperature(20 �C) using Hg[Co(SCN)4] as a calibrant; diamagneticcorrections were calculated from Pascal’s constants [29b].Thermogravimetric (TGA) studies were recorded on aSeteram Labsys TG-16 thermobalance. UV–Vis spectrawere recorded on an Agillent Model 8453 diode arrayspectrophotometer. Cyclic voltammograms (CV) were car-ried out using CV measurements with a Princeton AppliedResearch Model 2263 potentiostat controlled by an exter-nal PC. A three electrode system (BAS model solid cellstand) was used for CV measurements in DMSO and con-sisted of a 2 mm sized platinum disc electrode as the work-ing electrode, a platinum wire counter electrode and anAg/AgCl reference electrode. The reference electrode wasseparated from the bulk solution by a fritted-glass bridgefilled with the solvent/supporting electrolyte mixture. Theferrocene/ferrocenium couple (Fc/Fc+) was used as aninternal standard, but all potentials in the paper are refer-enced to the Ag/AgCl reference electrode. Solutions con-taining the complexes were deoxygenated by a stream ofhigh purity nitrogen for at least 5 min before running theexperiment, and the solution was protected from air by ablanket of nitrogen during the experiment. Controlledpotential electrolysis (CPE) was performed with a Prince-ton Applied Research Model 2263 potentiostat/galvano-stat. A BAS model electrolysis cell with a fritted glass toseparate the cathodic and anodic portions of the cell wasused for bulk electrolysis. The sample and solvent wereplaced into the electrolysis cell under nitrogen. UV–Visspectroelectrochemical experiments were performed witha home-built thin-layer cell that utilized a light transparentplatinum gauze working electrode [30]. Potentials wereapplied and monitored with a Princeton Applied ResearchModel 2263 potentiostat. Time-resolved UV–Vis spectrawere recorded on an Agillent Model 8453 diode arrayspectrophotometer.

A. Kilic et al. / Polyhedron 26 (2007) 4009–4018 4011

2.2. Synthesis of the ligands L1 and L2

The ligands N-N 0-bis(3,5-di-tert-butylsalicylaldimine)-1,4-cyclohexane bis(methylamine) (L1) and N-N 0-bis(3,5-di-tert-butylsalicylaldimine)-1,8-diamino-3,6-dioxaoctane(L2) were synthesized by the reaction of 6 mmol 3,5-di-tert-butyl-2-hidroxybenzaldehyde in 20 ml absolute ethanolwith 3 mmol 1,4-cyclohexane bis(methylamine) and3 mmol 1,8-diamino-3,6-dioxaoctane in 10 ml ethanol,respectively. Also, 3–4 drops of acetic acid were added.The mixtures were refluxed for 3 h, followed by coolingto room temperature. The crystals were filtered in vacuum,then the products were recrystallized from ethanol. Theproducts are soluble in common solvents such as CHCl3,DMF and DMSO.

For ligand L1: Color: yellow; mp: 194 �C; yield (%): 78;Anal. Calc. for C38H58N2O2: C, 79.44; H, 10.10; N, 4.88.Found: C, 79.31; H, 10.15; N, 4.74%. 1H NMR (CDCl3,TMS, d ppm): 13.78 [(OH), s, 2H], 8.30 and 8.33 [(CH@N),d, 2H, (J = 14.8 Hz)], 1.32 [(C–CH3), s, 18H], 1.45 [(C–CH3),s, 18H], 7.05 [(Ar-H), s, 2H], 7.35 [(Ar-H), s, 2H], 3.46 and3.57 [(N–CH2), d, 4H, (J = 5.6 Hz and 6.8 Hz)], 1.56–1.91[(Cyc-CH2), m, 8H], 1.08–1.13 [(Cyc-CH–), m, 2H]. 13CNMR (CDCl3, TMS, d ppm): C1,7 (29.73), C2,6 (31.82 and31.00), C3 (126.78), C4 (136.86), C5 (125.91 and 125.89), C8

(158.43), C9 (118.06 and 118.04), C10 (139.95 and 139.91),C11 (166.00 and 166.03), C12 (66.63 and 63.79), C13 (39.34and 36.70), C14 (35.28 and 34.35). IR (KBr pellets, cm�1):3410 m(OH), 2957, 2865 m(Aliph-H), 1633 m(C@N), 1175m(C–O). UV–Vis [in C2H5OH kmax/nm (log e), (*: shoulderpeak)]: 262 (3.82), 331 (3.15), 419* (2.32).

For ligand L2: Color: yellow; mp: 106 �C; yield (%): 71;Anal. Calc. for C36H56N2O4: C, 74.48; H, 9.65; N, 4.82.Found: C, 73.76; H, 10.03; N, 4.46%. 1H NMR (CDCl3,TMS, d ppm): 13.66 [(OH), s, 2H], 8.35 [(CH@N), s, 2H],7.02 [(Ar-H), s, 2H], 7.36 [(Ar-H), s, 2H], 3.71 and 3.74[(Ar-H), d, 8H], 3.63 [(N–CH2), s, 4H], 1.31 [(C–CH3), s,18H], 1.45 [(C–CH3), s, 18H]. IR (KBr pellets, cm�1):3413 m(OH), 2960, 2876 m(Aliph-H), 1633 m(C@N), 1173m(C–O), 1139 m(C–O–C). UV–Vis [in C2H5OH kmax/nm(log e), (*: shoulder peak)]: 251* (4.19), 262 (5.02), 329(4.02), 414* (2.54).

2.3. Synthesis of the metal complexes

2 mmol of the ligands L1 and L2 were dissolved in 40 mlabsolute methanol and 2 mmol Cu(Ac)2 Æ 2H2O and Co-(Ac)2 Æ 4H2O in 20 ml methanol were mixed. The mixtureswere heated at 50 �C with constant stirring until they wereconcentrated. The mixtures were evaporated to a volume of15–20 ml in vacuum and left to cool to room temperature.The compounds were precipitated after adding 5 ml etha-nol. The products were filtered in vacuum and washed witha small amount of methanol and water. The products wererecrystallized from methanol. All the complexes are stableat room temperature and soluble in solvents such asCHCl3, DMF and DMSO.

[Co(L1)]: Color: dark yellow; mp: 214 �C; yield (%):67; KM: 1.2 X�1 cm2 mol�1; leff = 3.61 BM. Anal. Calc.for C38H56N2O2Co: C, 72.26; H, 9.19; N, 4.44. Found: C,72.58; H, 8.94; N, 4.38%. IR (KBr pellets, cm�1): 2957,2866 m(Aliph-H), 1619 m(C@N), 1167 m(C–O), 488 (M–O),544 (M–N). UV–Vis [in CHCl3 kmax/nm (log e), (*: shoulderpeak)]: 227* (3.70), 263 (4.56), 330* (4.11).

[Cu(L1)]: Color: brown; mp: 232 �C; yield (%): 70; KM:3.5 X�1 cm2 mol�1; leff = 1.89 BM. Anal. Calc. forC38H56N2O2Cu: C, 71.75; H, 8.81; N, 4.40. Found: C,71.98; H, 8.65; N, 4.32%. IR (KBr pellets, cm�1): 2956,2861 m(Aliph-H), 1618 m(C@N), 1170 m(C–O), 493 (M–O),537 (M–N). UV–Vis [in CHCl3 kmax/nm (log e)]: 286(4.25), 328 (4.23), 380 (4.17).

[Co(L2)]: Color: dark yellow; mp: 123 �C; yield (%):65; KM: 1.8 X�1 cm2 mol�1; leff = 3.76 BM. Anal. Calc.for C36H54N2O4Co: C, 67.76; H, 8.47; N, 4.39. Found: C,67.13; H, 8.82; N, 4.01%. IR (KBr pellets, cm�1): 2955,2866 m(Aliph-H), 1615 m(C@N), 1171 m(C–O), 1124 m(C–O–C), 487 and 465 (M–O), 546 (M–N). UV–Vis [in CHCl3kmax/nm (log e), (*: shoulder peak)]: 233* (5.32), 248 (4.81),375* (4.24), 876* (1.66).

[Cu(L2)]: Color: black; mp: 195 �C; yield (%): 68; KM:3.9 X�1 cm2 mol�1; leff = 1.95 BM. Anal. Calc. forC36H54N2O4Cu: C, 67.34; H, 8.41; N, 4.36. Found: C,68.02; H, 8.76; N, 4.12%. IR (KBr pellets, cm�1): 2953,2865 m(Aliph-H), 1625 m(C@N), 1173 m(C–O), 1112 m(C–O–C), 490 and 466 (M–O), 534 (M–N). UV–Vis [in CHCl3kmax/nm (log e), (*: shoulder peak)]: 276* (5.24), 321 (4.87),364* (2.27), 443* (2.29), 665 (2.28)

3. Results and discussion

3.1. Synthesis

The reaction for the synthesis of the ligands is given inScheme 1. The first step is the synthesis of N-N 0-bis(3,5-di-tert-butylsalicylaldimine)-1,4-cyclohexane bis(methyl-amine) (L1) and N-N 0-bis(3,5-di-tert-butylsalicylaldimine)-1,8-diamino-3,6-dioxaoctane (L2) from the reaction of3,5-di-tert-butyl-2-hidroxybenzaldehyde with 1,4-cyclohex-ane bis(methylamine) and 1,8-diamino-3,6-dioxaoctane,respectively. The second step is the synthesis of [CuL1],[CuL2] and [CoL1], [CoL2] complexes from the reactionof N-N 0-bis(3,5-di-tert-butylsalicylaldimine)-1,4-cyclohex-ane bis(methylamine) (L1) and N-N 0-bis(3,5-di-tert-butyl-salicylaldimine)-1,8-diamino-3,6-dioxaoctane (L2) with thecorresponding metal salts (Scheme 2).

The metal to ligand ratios of the Co(II) and Cu(II) com-plexes were found to be 1:1. Infrared spectra of ligands andtheir complexes were recorded in KBr pellets from 4000 to400 cm�1. The IR spectra of the ligands are characterizedby the appearance of a band at 3410 cm�1 for L1 and3413 cm�1 for L2 due to the m(O–H) groups [31]. In thespectra of the Cu(II) and Co(II) complexes, these bandsdisappear. The vibrations of the azomethine groups ofthe free ligands are observed at 1633 cm�1. In the case of

+ 60 oC, CH3OH

-2H2ONH2 R NH2

(H3C)3C

(H3C)3C OH

CH

O

2R

N=CH

N=CHHO

HO

C(CH3)3

C(CH 3)3

C(CH 3)3

C(CH 3)312

34

56 78

9

10

11

R: O

Oor

(L1) (L2)

Scheme 1. Synthetic route to the ligands.

R

N=CH

N=CHHO

HO

C

CH3

CH3

CH3

:

CH3OH

50 oCR

N=CH

N=CHO

OM

1) Cu(Ac) 2.2H2O2) Co(Ac) 2.4H2O

Ligands (L) M= Co(II) or Cu(II)

R: O

Oor

(L1) (L2)

Scheme 2. Synthetic route to the complexes.

4012 A. Kilic et al. / Polyhedron 26 (2007) 4009–4018

the complexes, these bands are shifted to the lower frequen-cies, indicating that the nitrogen atom of the azomethinegroups is coordinated to the metal ion, as expected [32].The bands at 1175 cm�1 for L1 and 1173 cm�1 for L2 areascribed to the phenolic C–O stretching vibrations. Thesebands are shifted to lower frequencies due to O-metal coor-dination [33]. Additionally, a peak observed at 1139 cm�1

in the free ligand (L2) has been assigned to the ethericC–O–C stretching. On complexation, this band is shifted

to a lower frequency range 1124–1112 cm�1, indicatingcoordination through the etheric oxygen. The coordinationof the azomethine nitrogen is confirmed with the presenceof a new band between 534 and 546 cm�1, assignable tom(M–N) for the Co(II) and Cu(II) complexes [34]. Thecoordination of the azomethine oxygen group is confirmedwith the presence of a new band at between 465 and493 cm�1, assignable to m(M–O) for the Co(II) and Cu(II)complexes.

A. Kilic et al. / Polyhedron 26 (2007) 4009–4018 4013

The 1H NMR and 13C NMR spectra of the ligands L1

and L2 were recorded in CDCl3 solution at 400 MHz. The1H NMR spectrum of L2 and the 13C NMR spectrum ofL1 are given in Fig. 1. The 1H NMR spectra of the ligandsL1 and L2 showed one singlet peak at d = 13.78 ppm andd = 13.66 ppm due to –OH protons, respectively. Also,the L1 and L2 ligands indicate the presence of two equiva-lent azomethine groups. Due to the equivalent chemicalenvironments, two signals for L1 at 8.30 and 8.33 ppm (d,2H) and one signal for L2 at 8.35 ppm (s, 2H) are recordedfor the azomethine protons [35]. However, the protons ofthe tert-butyl groups of the ligands exhibit two sharp singletpeaks between 1.31 and 1.45 ppm, indicating that the tert-butyl protons of these compounds are magnetically non-equivalent [36,37]. On the other hand, the 1H and 13CNMR spectra of the ligand L1 showed signals at 166.07and 166.03 ppm belonging to C11, 66.63 and 63.79 ppmbelonging to C12, CH@N (8.30 and 8.33 ppm) and N–CH2 (3.53–3.57 and 3.45–3.46 ppm). These carbon and pro-ton resonances have a double resonance indicating thatligand (L1) has cis–trans isomerism [38]. The isomer ratiowas found to be 60% cis-isomer and 40% trans-isomer fromthe 1H and 13C NMR data. The NMR data of the ligands

Fig. 1. The 1H NMR spectrum of L2 (a); the 13C NMR spectrum of L1

(b).

and their complexes are given in Section 2. Since the Co(II)and Cu(II) complexes are paramagnetic nature, their NMRspectra could not be obtained.

The electronic spectra of the ligands and their metalcomplexes were recorded in C2H5OH for the ligands andin CHCI3 for the complexes. The bands below 331 nmare attributable to intramolecular p! p* and n! p* tran-sitions. The very intense bands at low wavelengths havebeen assigned to charge transfer transitions [39,40], formetal complexes with aromatic bridges these bands occurat longer wavelengths, as expected from the higher aroma-ticity of the ligands which eases delocalization of electrondensity.

3.2. Magnetic measurements

Room temperature magnetic moments of the Cu(II)complexes are between 1.89 and 1.95 BM, which is typicalfor mononuclear compounds of Cu(II) with a S = 1/2 spin-state and does not indicate antiferromagnetic coupling ofthe spins at this temperature. The room temperature mag-netic moment of between 3.61 and 3.76 BM determined forthe Co(II) complexes are close to the spin-only magneticmoments (l = 3.87 BM) for three unpaired electrons.

3.3. Thermal analysis

The thermal stability of the ligands and the Co(II) andCu(II) metal complexes was investigated by TGA. TheTGA curves were obtained at a heating rate of 5 and10 �C min�1 under a nitrogen atmosphere over the 20–900 �C range, and they show that L1 and L2 and the Cu(II)and Co(II) complexes are thermally stable up to 198 for L1,164 for L2, 240 for [Cu(L1)], 214 for [Co(L1)], 174 for[Cu(L2)] and 189 for [Co(L2)] �C. As a results of the ther-mal analysis, the most thermally stable complex is[Cu(L1)], whereas the least stable complex is [Cu(L2)] [35c].

3.4. Molar conductivity

The molar conductivities (KM) of the Cu(II) and Co(II)metal complexes in DMF (dimethyl formamide) at 10�3 Mwere found to be in the range 1.2–3.9 X�1 cm2 mol�1. Theselow values indicate that all of the metal complexes are non-electrolytes due to no counter ions in the proposed structuresof the Cu(II) and Co(II) metal complexes (Scheme 2)[35b,35c]. The difference in the number of protons in theCu(II) and Co(II) metal complexes probably points to thedifference in size of the +2 oxidation state of these metal ions,with the different polarizing ability due to the different sizes.Thus, the electrical conductivities of the Cu(II) complexesare higher than the Co(II) metal complexes [41,42].

3.5. Electrochemistry

The electrochemical behavior of [Cu(L1)], [Cu(L2)] and[Co(L1)], [Co(L2)] was investigated using the cyclic

Table 1The electrochemical potentials of the copper and cobalt complexes

Complexes Electrolyte Solvent M(II)/M(I) Epcb DEa (V) M(II)/M(III) Epa

b DEa (V)

[Cu(L1)] TBAP CH2Cl2 1.40[Cu(L2)] TBAP CH2Cl2 �1.07 358 1.09[Co(L1)] TBAP CH2Cl2 0.859[Co(L2)] TBAP CH2Cl2 0.348 248

a DEp = (Epa � Epc).b Anodic peak potential and cathodic peak potential for the irreversible process.

4014 A. Kilic et al. / Polyhedron 26 (2007) 4009–4018

voltammetric (CV) technique in CH2Cl2 solution contain-ing 0.1 M TBAP. The data obtained in this work are listedin Table 1. Fig. 2 shows the CVs of [Cu(L1)] and [Cu(L2)] at0.100 V s�1 scan rate. As can be seen from Fig. 2a, [Cu(L1)]displayed one oxidation wave in CH2Cl2 containing theAg/AgCl electrode system. The anodic peak potential forthe oxidation process displayed at Epa = 1.40 V versusAg/AgCl. The oxidation process has a large value for thecathodic to anodic peak separation, and deviates fromunity for the peak cathodic and peak anodic current ratios,ipc/ipa, at 0.100 V s�1 scan rate. The results show thatthe redox process is irreversible in the scan range of

Cur

rent

/ μA

Potential / V vs. Ag/AgCl

Cur

rent

/ μA

Potential / V vs. Ag/AgCl-1.5 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0

0

5

10

15

Ic

Ia

-1.5 -1.0 -0.5 0.0 0.5 1.0 1.5-5

0

5

10

II'Ic

IIa

Ia

a

b

Fig. 2. Cyclic voltammograms of [Cu(L1)] (a); [Cu(L2)] (b) in CH2Cl2solution containing 0.1 M TBAP. Scan rates = 0.100 V s�1.

0.025–0.250 V s�1 in CH2Cl2 solution, and can be indica-tive of a decomposition reaction occurring under the CVconditions. As can be seen from Fig. 2b, [Cu(L2)] exhibiteda totally irreversible oxidation process without a corre-sponding anodic wave at 0.100 V s�1 scan rate, and its ano-dic potential displayed at Epa = 1.09 V versus Ag/AgCl. Onreverse scan, a cathodic wave (II 0) appeared at 0.200 V fol-lowing the oxidation potential (II), which was assigned toan EC process. [Cu(L2)] also showed one reduction process(Ic) at Epc = �1.07 V. This reduction wave is an irreversibleprocess because of the large value of the cathodic to theanodic peak separation (DE = 358 V at 0.100 V s�1 scanrate), even if it has the corresponding anodic wave (Ia).The results reveal that the [Cu(L1)] and [Cu(L2)] complexesexhibit different electrochemical behavior under the sameexperimental conditions. It is well known from the litera-ture [43] that redox potentials of Schiff-base complexesare markedly dependent on the ligand, solvent and geome-try of the complexes. Electron-withdrawing substituents onthe ligands shift the redox potential positively, whereaselectron-donating groups have the opposite effect. On theother hand, the oxidation potentials are also dependenton the imine bridge of the Schiff-base, with the more posi-tive potentials observed for complexes with aromaticbridges. Two possible explanations can be invoked toexplain this behavior. (i) An increase in unsaturation thatwould ease electron-density delocalization from the metalto the ligand, and (ii) higher rigidity of the ligand, impartedby the aromatic bridge, that would hinder a contraction ofthe hole cavity. Moreover, the redox potentials show acathodic shift with an increase in the donor capacity ofthe solvent. For example, oxidation of square-planarmetal(II) complexes to metal(III) takes place with a con-comitant increase in coordination number, as these lattercomplexes are normally six-coordinate. The cathodic shiftin the redox potentials ongoing from DMF to DMSO isattributed to the stronger axial bonds that are establishedbetween the metal(III) center and the molecules of the bet-ter donor, and that help to stabilize the high oxidation stateof the metal [43a]. The size of the coordination cavity in thecomplexes, and the geometric requirements and the size ofthe metal ions in different oxidation states also affect theredox potential [44]. For example, a tetrahedral structureof the Schiff-base complex enhances the ease of reductionof Cu(II) to Cu(I) compared to a planar coordinationabout the copper atom [43e]. In our case, the experiments

0.0 0.2 0.4 0.6 0.8 1.0

0

1

2Ia

-1.5 -1.0 -0.5 0.0 0.5 1.0

0

1

Ic

Ia

Cur

rent

/ μA

Potential / V vs. Ag/AgCl

Cur

rent

/ μA

Potential / V vs. Ag/AgCl

a

b

Fig. 3. Cyclic voltammograms of [Co(L1)] (a); [Co(L2)] (b) in CH2Cl2solution containing 0.1 M TBAP. Scan rates = 0.100 V s�1.

A. Kilic et al. / Polyhedron 26 (2007) 4009–4018 4015

were carried out in CH2Cl2 solution so any solvent effectcan be eliminated on comparing the electrochemical behav-ior. Taking into account substituent effects for the two cop-per Schiff-base complexes, both cyclohexane anddioxaoctane bridged on the imine group of the Schiff-basecan be considered as weak electron-donating groups, thus alarge potential shift could not be expected for these coppercomplexes. However, the structural differences of the[Cu(L1)] and [Cu(L2)] complexes should be taken intoaccount for the different electrochemical behavior of thecomplexes under the same experimental conditions.Because the CV of [Cu(L2)] showed a lower oxidationpotential compared to [Cu(L1)], it was seen that the oxida-tion wave (II) of [Cu(L2)] was accompanied by an EC pro-cess (II 0), which was not observed for [Cu(L1)]. Also,[Cu(L2)] gave a reduction process but [Cu(L1)] did not. Ifthe anodic scan was limited to the potential E = 0.6 V,the EC process was not observed. These results indicatethat the Cu(II) ion in [Cu(L2)] is coordinated by N2O4

donor sites where two complementary lone pairs comefrom dioxaoctane binding on the imine group of theSchiff-base responsible for this binding mode of the ligand.On the other hand, [Cu(L1)] presents a square-planar struc-ture with four donor atoms, as observed for several Schiff-base complexes with N2O2 donor groups. In the case of[Cu(L2)], the oxidation of Cu(II) (d9) to Cu(III) (d8)involves a drastic decrease in the metal radius and a confor-mation change from octahedral to square-planar. Thisstructural reorganization of the oxidized complex resultedin the observation of the EC process. This process (II 0), fol-lowing the oxidation wave (II) in the CV of the complex,exhibited re-reduction of the oxidized copper ion(Cu(III)! Cu(II)) where the complex changed into its ori-ginal form. The reduction process was observed for octahe-dral [Cu(L2)] complex but not for the square-planar[Cu(L1)] complex. The reduction of Cu(II) (d9) to Cu(I)(d10) involves a drastic increase in the metal radius, so[Cu(L2)] favors its original geometry. Therefore, no ECprocess was observed following the reduction process(Ic). The fact that [Cu(L2)] exhibits a reduction processcan be attributed to the different geometry of the complexcompared to [Cu(L1)].

The electrochemical behavior of the [Co(L1)] and[Co(L2)] complexes is presented in Fig. 3. On the positivescan, [Co(L1)] gave one oxidation wave without a corre-sponding cathodic wave (Ia) (Fig. 2a). The process istotally irreversible and its anodic peak potential is atEpa = 0.859 V versus Ag/AgCl. On the other hand,[Co(L2)] exhibited one oxidation wave with a correspond-ing cathodic wave. This oxidation process is diffusion con-trolled with the anodic current function (Ipa/v1/2)independent of the scan rate (v) over the scan range0.025–0.250 V s�1. The peak potentials separation isDEp = 248 mV and the ratio of cathodic to anodic peakcurrent deviates from unity. So the oxidation process isirreversible at all sweep rates in the range 0.025–0.250 V s�1. Both oxidation processes for [Co(L1)] and

[Co(L2)] are based on the cobalt center and are assign tothe Co(II)! Co(III) process. The electrochemical behav-ior of [Co(L2)] differs from that of [Co(L1)] under the sameexperimental conditions due to their different structures, asobserved for the copper complexes. The oxidation potentialof [Co(L2)] cathodically shifts compared with that of[Co(L1)]. An EC process following the oxidation wave(II) in the CV of the [Co(L2)] was not observed, probablydue to favoring an octahedral structure in the case ofCo(III) (d6), which indicates no structural change duringthe oxidation process.

3.6. UV–Vis spectra of the electro-oxidized and

electro-reduced complexes

The spectroelectrochemical behavior of the copper andcobalt complexes was investigated using an in situ spectro-electrochemical technique including chronoamperometryand UV–Vis spectroscopy in CH2Cl2 solution containing0.2 M TBAP. The UV–Vis spectral changes for the reducedand oxidized species of the corresponding complexes wereobtained in a thin-layer cell with the applied potentials.

300 400 500 6000.0

0.1

0.2

0.3

0.4

320 nm

416 nm

372 nm

Abs

orba

nce

Wavelength / nm

Fig. 5. Time-resolved UV–Vis spectral changes of [Cu(L2)] during thereduction at Eapp = �1.52 V in CH2Cl2 solution containing 0.2 M TBAP.

4016 A. Kilic et al. / Polyhedron 26 (2007) 4009–4018

The absorption spectra of the neutral complexes and theirelectrochemically generated species are given in Figs. 4–6.The convenient applied potentials values for the in situ

spectroelectrochemical experiment were determined foreach process by taking CVs of the complexes in the thin-layer cell.

Fig. 4a shows the changes of the electronic spectrum of[Cu(L1)] that is accompanied by the oxidation process inthe thin-layer cell. Well-defined isosbestic points areobserved at k = 321 and 328 nm, confirming that the elec-trode reaction proceeds in a quantitative fashion and there-fore the absence of any coupled chemistry. The band at327 nm, assigned to the Cu(II) dx2�y2 -to-phenolate (p)LMCT transition for the neutral copper complex[Cu(L1)], changed into two new bands. The first, withhigher intensity, shifted to a lower energy (k = 362 nm)and the second, also with higher intensity, appeared at ahigher energy (293 nm) as a result of removing one electronfrom the copper(II) ion in the complex. The one electronoxidation results in a Cu(II) (d9)! Cu(III) (d8) process

Abs

orba

nce

Wavelength / nm

Abs

orba

nce

Wavelength / nm

300 400 500 6000.0

0.5

1.0

1.5

-300 -200 -100 0 100 200 300-2

0

2

Cu

rren

t / m

A

Time / s

327 nm362 nm

293 nm

300 400 500 600

0.2

0.4

0.6

0.8

-0.5 0.0 0.5 1.0 1.5

0

10

20

E = 1.27 V

Cu

rren

t / μ

A

Potential / V vs. Ag/AgCl

303 nm

372 nm

a

b

Fig. 4. Time-resolved UV–Vis spectral changes of [Cu(L1)] (a); [Cu(L2)](b) during the oxidation at Eapp = 1.27 V in CH2Cl2 solution containing0.2 M TBAP. For (a), the inset figure shows chronoamperometry of[Cu(L1)]. For (b), the inset figure shows the thin-layer CV of [Cu(L2)]under the same experimental conditions.

which makes the electron transition from the ligand tothe metal easier, as expected. So the LMCT was observedat a lower energy compared with that of the neutral com-

300 350 400 450 5000.0

0.5

1.0

1.5

410 nm

375 nm333 nm

300 350 400 450 500 550 6000.0

0.2

0.4

0.6

-300 -200 -100 0 100 200 300-0.4

-0.2

0.0

0.2

0.4

Cur

rent

/ m

A

Time / s407 nm

375 nm

Abs

orba

nce

Wavelength / nm

Abs

orba

nce

Wavelength / nm

a

b

Fig. 6. Time-resolved UV–Vis spectral changes of [Co(L1)] (a); [Co(L2)](b) during oxidation at Eapp = 1.00 V for the first and Eapp = 0.75 V forthe latter in CH2Cl2 solution containing 0.2 M TBAP. For (b), the insetfigure shows chronoamperometry of [Co(L2)].

A. Kilic et al. / Polyhedron 26 (2007) 4009–4018 4017

plex. The second band, observed at 293 nm and corre-sponds to an intramolecular p! p* transition, dominatedwhen the LMCT transition shifted to lower energy. As seenfrom Fig. 4a and b, the final spectra (green solid lines1) ofthe oxidized species for the [Cu(L1)] and [Cu(L2)] com-plexes exhibited similar features. The corresponding mainbands of the oxidized species of [Cu(L1)] and [Cu(L2)]appeared at 293 and 362 nm, and 303 and 372 nm, respec-tively. Comparing the main bands, a small red shift for theoxidized [Cu(L2)] complex is seen, probably due to its dif-ferent environment. However, one can also observe anabsorption band at about 450 nm in the final spectrum ofthe oxidized [Cu(L1)], but, as seen Fig. 4a, this bandappears with a very low intensity compared the mainbands, which is probably due to the result of the differentsolubility degree of the resultant products in the thin layercell. It is evident that the bands of [Cu(L2)] are seen with alower intensity compared with those of [Cu(L1)] under thesame experimental conditions. Hence, this weak bandcould not be attributed to structural changes for theoxidized product of [Cu(L1)]. These spectral changes aresimilar to those previously observed for some copperSchiff-base complexes [28]. Although the oxidation processwas observed to be electrochemically irreversible from theCV of [Cu(L1)], the original spectrum of the complex wasobtained upon re-reduction during the spectroelectrochem-ical measurements. The controlled potential coulometric(CPC) study indicated that the number of electrons trans-ferred for the forward and reverse electrochemical reac-tions of the complex was one for both the oxidation andre-reduction processes, based on the Cu(II)/Cu(III) cou-ples. Thus, the oxidized [Cu(L1)] complex remained chem-ically stable throughout the experiment indicating theoxidized [Cu(L1)] complex remains in its original geometry.

The UV–Vis spectral changes upon oxidation of[Cu(L2)] are shown in Fig. 4b where the inset figure exhibitsthe CV of [Cu(L2)] in a CH2Cl2 solution in the thin-layercell, containing 0.2 M TBAP. The neutral complex showedtwo intense bands at 320 and 372 nm, which can be attrib-uted to an intramolecular p! p* transition and Cu(II)dx2�y2 -to-phenolate (p) LMCT transition, respectively [45].The LMCT transition of [Cu(L2)] is shifted to a lowerenergy compared with that of [Cu(L1)] as a result of the dif-ferent binding mode of the ligand of [Cu(L2)] where twoadditional oxygen lone pairs coming from dioxaoctanecoordinate to the copper ion. As can be seen fromFig. 4b, the isosbestic points of the UV–Vis spectralchanges were found at 329 nm and 359 nm, which clearlyverify the formation of the mono-oxidized species of[Cu(L2)]. The band intensity corresponding to the 372 nmpeak decreased, but a shift to a lower energy was notobserved. The final spectrum of the oxidized species of[Cu(L2)] is similar to that of [Cu(L1)] with the only differ-

1 For interpretation of color in Fig. 4, the reader is referred to the webversion of this article.

ence being that the absorption bands of the oxidized[Cu(L2)] are shifted to a lower energy compared with thoseof [Cu(L1)] because of their different coordination environ-ment. Finally, the oxidized species of the two copper com-plexes present a square-planar structure. The originalspectrum of the neutral complex [Cu(L2)] was not recov-ered upon re-reduction at the applied potential Eapp =0.6 V, but could be recovered upon re-reduction at theapplied potential Eapp = �0.3 V. These result indicate thatthe geometry of [Cu(L2)] changed into square-planar afterthe complex was totally oxidized and the neutral complexwas only recovered following the EC process observed inthe CV of [Cu(L2)].

The well-defined UV–Vis spectral changes observed dur-ing the reduction of [Cu(L2)] are shown in Fig. 5. After thereduction process, one electron was included in the metalsystem, thus a Cu(II) (d9)! Cu(I) (d10) process took place.As expected, the ligand to metal charge transfer (LMCT) at372 nm decreases and a relatively strong peak at 416 nmdevelops. The absorption maximum at 416 nm is attributedto an MLCT transition in the Cu(I) complexes. Thesechanges are in agreement with the expected results for ametal-centered reduction, Cu(II)! Cu(I) [27,28].

The UV–Vis spectral changes upon oxidation of[Co(L1)] and [Co(L2)] are shown in Fig. 6. The electronicspectra of the Co(II) complexes showed intense absorptionbands at 330 and 375 nm that are assigned to a p! p*

transition associated with the imine linkage. For the twocobalt complexes, the bands corresponding to the p! p*

transition disappeared and new bands with small red shiftsand of lower intensity were observed during the oxidationprocess. These new bands are attributed to the LMCTtransition. Isosbestic points are observed at 330 and408 nm for [Co(L1)] and 330 and 408 nm for [Co(L2)].The final spectra of the oxidized species of the cobalt com-plexes are similar to each other and to those previouslyreported for the oxidations of Schiff base cobalt complexes[28]. Re-reduction at a negative potential of the oxidizedspecies of [Co(L2)] generates the neutral complex [Co(L2)]indicating that the oxidized species is chemically stablethroughout the spectroelectrochemical timescale, probablydue to the six coordination of the [Co(L2)] complex, whichis able to stabilize the oxidized species compared to the tet-rahedral [Co(L1)] complex for which the original spectrumcould not be recovered upon re-reduction [43a].

4. Conclusion

New Schiff-base copper and cobalt complexes, [Cu(L1)],[Cu(L2)], and [Co(L1)], [Co(L2)] (where L1 = N-N 0-bis(3,5-di-tert-butylsalicylaldimine)-1,4-cyclohexane bis(methyl-amine) and L2 = N-N 0-bis(3,5-di-tert-butylsalicylaldi-mine)-1,8-diamino-3,6-dioxaoctane), were synthesized andtheir electro-spectrochemical properties were investigatedusing cyclic voltammetric and thin-layer spectroelectro-chemical techniques in a dichloromethane solution(CH2Cl2). The combination of the electrochemical and

4018 A. Kilic et al. / Polyhedron 26 (2007) 4009–4018

spectroelectrochemical methods provided a powerful toolto reveal the complementary nature of the molecular struc-ture and electron-transfer reactions of the oxidized andreduced species of the new electroactive Schiff-base com-plexes. The comparative electro-spectrochemical studiesrevealed that the [Cu(L1)] and [Co(L1)] complexes showedfour coordination with the ligand (L1), having two N andtwo O donor sites, while the [Cu(L2)] and [Co(L2)] com-plexes resulted in six coordination with the ligand (L2) car-rying two N and four O donor sites. The electro-oxidizedspecies of the [Cu(L1)] and [Co(L2)] complexes remainedstable on the spectroelectrochemical timescale, althoughthe complexes exhibited an irreversible oxidation processesin their CV measurements. Finally, the CVs and thewell-defined spectral changes observed with the appliedpotentials in the thin-layer cell spectral data support thecorresponding structures of the copper and cobalt Schiff-base complexes.

Acknowledgements

This work has been supported, in part, by the ResearchFund of Harran University (Sanliurfa, Turkey). The sup-port of the State Planning Organization (DPT, ProjectNo. 90177) is gratefully acknowledged. This work has alsobeen supported, in part, by the Turkish Academy of Sci-ences in the framework of the Young Scientist Award Pro-gram (TUBA-GEB_IP).

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