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Gentamicin complexes of cobalt(II),nickel(II), cadmium(II) and tin(II)Parashuram Mishra a & Rakesh K. Sharma aa Department of Chemistry, University of Delhi, Delhi, 110007,IndiaVersion of record first published: 04 Nov 2008.

To cite this article: Parashuram Mishra & Rakesh K. Sharma (2008): Gentamicin complexes ofcobalt(II), nickel(II), cadmium(II) and tin(II), Main Group Chemistry, 7:2, 83-95

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Gentamicin complexes of cobalt(II), nickel(II),

cadmium(II) and tin(II)

Parashuram Mishra and Rakesh K. Sharma*

Department of Chemistry, University of Delhi, Delhi 110007, India

(Received 13 May 2008)

Gentamicin is an aminoglycoside antibiotic with known adverse side effects to thekidney and the inner ear. It is proposed that these effects result from complex formationof gentamicin with Co (1), Ni (2), Cd (3) and Sn (4). Metal coordination to gentamicinwas studied by various spectroscopic techniques. The crystal system, lattice parameters,unit cell, particle size and volume were determined by XRPD. The crystallographicdata of complex (2) are: Monoclinic, space groupP4/m, a ¼ 16.67920A, b ¼ 12.71476A,c ¼ 5.2673A, V ¼ 1112.00A3, a ¼ 90, b ¼ 95.21, g ¼ 908 and for complex (4): crystalsystem, Triclinic, space group, P1, a ¼ 8.290549, b ¼ 9.619387, c ¼ 16.12853,a ¼ 63.75208, b ¼ 53.3017, g ¼ 62.28692, V ¼ 888.68. The molecular structures ofthe complexes were optimized by MM2 calculations which indicated a square planararrangement of gentamicin around the metal ions.

Keywords: aminoglycoside; gentamicin; spectra; metal binding; molecular modelling

Introduction

Gentamicin C (Figure 1) is one of the aminoglycoside antibiotics having chemical name2-(4,6-diamino-3-(5-amino-4-(1-(methylamino)ethyl)tetrahydro-2H-pyran-2-yloxy)-2-hydro-xycyclohexyloxy)-5-methyl-4-(methylamino)tetrahydro-2H-pyran-3,5-diol that is widelyused for the treatment of serious infections caused mainly by Gram-negative and someGram-positive bacteria [1]. Commercial gentamicin C contains three closely relatedanalogues, which differ in the degree of methylation of the purpurosamine unit.Aminoglycosides kill bacteria primarily by inhibiting the translation step in microbialprotein synthesis [2]. A major problem in therapy with aminoglycosides is their relativelyhigh toxicity to the kidney and the inner ear. Nevertheless, gentamicin is currently the firstchoice antibiotic in developing countries and is still widely used in industrialized countriesfor the treatment of serious bacterial infections. The adverse effects of aminoglycosides mayresult from complex formation with transition metal ions and the oxidative reactions thecomplexes subsequently promote [3–5]. Co-administration of transition metal chelatorsand free radical scavengers, as well as over-expression of superoxide dismutase in modelanimals, suppress aminoglycoside-induced ototoxicity [6–9]. Systematic in vitro studies ofiron interactions with gentamicin have led to a postulated mechanism of toxicity involvingfree radical formation by a Fe(II)/Fe(III)–gentamicin complex. Another body of evidence[10–12] has suggested that both pharmacological activity and toxicity of aminoglycoside

*Corresponding author. Email: prmmishra@redimail.com

Main Group Chemistry

Vol. 7, No. 2, June 2008, 83–95

ISSN 1024-1221 print/ISSN 1745-1167 online

� 2008 Taylor & Francis

DOI: 10.1080/10241220802197469

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antibiotics could be related to copper(II)-aminoglycosides complexes. Jezowska-Bojczuket al. extensively investigated chelation of copper(II) ions by gentamicin-relatedaminoglycoside antibiotics using potentiometry and a variety of spectroscopic techniques[10–13]. Kanamicin B, tobramicin, geneticin and amikacin strongly bind Cu(II) ions,forming monomeric complexes over a wide pH range. In naturally occurring aminoglyco-sides, the amino nitrogens and deprotonated alcoholic oxygens of the terminal aminosugarrings are involved in the coordination, forming five- and six-membered chelate rings aboutcentral ions [14–16]. Amikacin, a semisynthetic derivative of kanamicin A, having the 1-amino group on the 2-deoxystreptamine moiety modified by acylation with 4-amino-2-hydroxybutyric acid, exhibits different binding modes by involving the amidated nitrogenin coordination [17]. Further, Cu(II)-amikacin complexes catalyze hydrogen peroxidedisproportionation at pH 7.4 mediated by hydroxyl radicals and involving Cu(I)/Cu(II)and Cu(II)/Cu(III) redox pairs [18]. These complexes mediate oxidation of 20-deoxyguanosine to 7,8-dihydro-8-oxo-20-deoxyguanosine, double-stranded DNA clea-vage, and both hydrolytic and oxidative t-RNAPhe strand scission at a specific site in theanticodon loop [19]. Under these circumstances, copper(II) ions are proposed to beinvolved in aminoglycoside toxicity.

L.C. Yan and coworkers studied the interaction of other aminoglycosides, such askanamicin A and neomycin copper(II) complexes, with a variety of nucleic acids [20]. Inthe presence of H2O2 and ascorbic acid these complexes promote highly specific double-stranded DNA scission by binding to the minor groove of DNA and abstracting the 40-Chydrogen of the deoxyribose moiety [21]. Cu(II)-kanamicin A complexes are also moreeffective than free kanamicin A in cleaving RNA targets within bacterial cells, suggesting apotential application for copper(II)-aminoglycoside complexes as new metal-based drugs[21]. Furthermore, the literature on gentamicin as a ligand reveals that there is cleardisagreement about the metal–ligand binding sites in complexes of this ligand [22].

Figure 1. Gentamicin C.

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Therefore, in order to ascertain metal binding sites in gentamicin metal chelates, thesynthesis and spectroscopic characterization of complexes of gentamicin C (Figure 1) withthese transition/main group metal ions is provided in this investigation.

2. Experimental

2.1. Materials and methodology

Gentamicin was purchased from Merck. All the chemicals used in this study were ofanalytical grade and used as procured. Solvents used were of analytical grade and werepurified by standard procedures. The stoichiometric analyses (C, H and N) of thecomplexes were performed using Elementar vario EL III (Germany) model. Metalcontents were estimated on an AA-640-13 Shimadzu flame atomic absorptionspectrophotometer in solution prepared by decomposing the respective complex in hotconcentrated HNO3. Their IR spectra were recorded on Perkins–Elmer FTIR spectro-photometer in KBr and polyethylene pellets. The electronic spectra were recorded in wateron Beckman DU-64 spectrophotometer with quartz cells of 1 cm path length. 1H NMRspectra were recorded in CDCl3 solvent on a Bruker Advance 400 instrument. Rigakumodel 8150 thermoanalyzer (Thermaflex) was used for simultaneous recording of TG-DTA curves at a heating rate of 108C min71. For TG, the instrument was calibrated usingcalcium oxalate while for DTA, calibration was done using indium metal, both of whichwere supplied along with the instrument. A flat bed type aluminum crucible was used witha-alumina (99% pure) as the reference material for DTA. The activation energy andArrhenius constant of the degradation process was obtained by Coats and Redfernmethod. The XRD powder pattern were recorded on a vertical type Philips 1130/00 X-raydiffractometer, operated at 40 kV and 50 Ma generator using the Cuka line at 1.54056 Aas the radiation sources. Sample was scanned between 58 and 708 (2y) at 258C. Thecrystallographic data were analyzed by using the CRYSFIRE –2000 powder indexingsoftware package and the space group was found by the GSAS program. Debye – Schererrelation with the help of 100% peak width determined the particle size. The density wasdetermined by Archimedes method.

2.2. 3D – Molecular modeling

3D molecular modeling of the proposed structure of the complexes was performedusing CsChem3D program package. The correct stereochemistry was assured throughthe manipulation and modification of the molecular coordinates to obtain reasonablelow energy molecular geometries. The potential energy of the molecule was the sum ofthe following terms: E ¼ Estr þ Eang þ Etor þ Evdw þ Eoop þ Eele. Where all Esrepresent the energy values corresponding to the given types of interaction. Thesubscripts str, ang, tor, vdw, oop and ele denote bond stretching, angle bonding, torsiondeformation, van der waals interactions, out of plain bending and electronic interaction,respectively.

2.3. Synthesis of complexes

To a methanolic solution of different metal chlorides (0.5 mmol) in a separate flask wasadded a methanolic solution of the gentamicin (0.5 mmol). The solution was stirred for6 h, after which the volume was reduced on a warm water bath. The product obtained waswashed with a small amount of methanol and air-dried. The above product was

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redissolved in excess warm methanol, and clear solution was left undisturbed for weeks togive beautiful crystals of the complexes.

3. Results and discussion

Satisfactory results of elemental analysis (Table 1) and spectral studies (Table 2) revealedthat the complexes were of good purity. Various attempts to obtain the single crystals haveso far been unsuccessful. X-ray diffraction studies indicate crystalline nature of the metalcomplexes. The complexes were soluble in polar solvents.

3.1. Vibrational spectra

Gentamicin molecule exhibits absorptions 1050,1750, 2934 cm71. These bands are verymetal complexes indicating non-involvement of the oxygen atoms of hydroxyl group incoordination with the metal ions [19]. The stretching frequencies of gentamicin hydroxylgive bands at 3368 and 3434 cm71 with a shoulder at about 3550 cm71. These bandsappear in the complexes as strong band absorption in the region 3420–3445 cm71. Thesebands appear for the new complex at the same wave number, ruling out the participationof hydroxyl oxygen in the coordination. These results confirm that complexation occurredand suggest that the oxygen of the hydroxyl group is involved in the coordination sphere[16]. The vibrational bands due to rocking and wagging modes of water and metal–oxygenstretching modes are observed in the 800–350 cm71 region for all the complexes may beattributed to coordinated water [20]. This can be confirmed with the help of thermograms.A new band in the 615–300 cm71 regions in the spectra of the complexes is assignableto n (M–O).

Table 1. Color, reaction yield and elemental analysis of complexes.

ComplexEmpiricalformula Color

Yield(%)

Analysis: found (calculated) (%)

C H N M

[Co(L)(2H2O)] (1) C21H45N5O9Co Pink 80 44.21(42.20)

7.95(7.91)

12.27(12.67)

10.33(10.35)

[Ni(L)H2O]H2O (2) C21H45N5O9Ni Green 81 44.24(44.42)

7.98(7.95)

12.29(12.28)

10.30(10.29)

[Cd(L)H2O] (3) C21H45N5O9 Cd White 80 40.45(40.42)

7.28(7.27)

11.28(11.21)

18.11(18.01)

[Sn(L) H2O] (4) C21H45N5O9Sn Brown 75 40.02(40.02)

7.21(7.20)

11.11(11.11)

18.81(18.83)

L, Gentamicin.

Table 2. IR spectral data (cm71) of the metal complexes.

Frequency nN–H OH OH NH2 NH2 M–O

[Co(L)H2O] (1) 3424 (s,b) 1634 (m) 1515 (s) 1220 (m) 695 (s) 350 (m)[Ni(L)H2O] (2) 3429 (s,b) 1638 (m) 1510 (s) 1210 (m) 690 (s) 465 (s)[Cd(L)H2O] (3) 3445 (s,b) 1644 (m) 1508 (m) 1219 (w) 692 (s) 372 (m)[Sn(L)H2O] (4) 3421 (s,b) 1625 (m) 1489 (s) 1215 (w) 695 (s) 365 (s)

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3.2. 1H NMR spectra

The 1H NMR spectra in a DMSO–d6 solvent of the ligand and Sn (II)–gentamicincomplex show well-resolved signals (Figure 4: 1H NMR spectrum of complex I) The N–Hprotons of amine, which would have undergone very rapid exchange with the solvent,appear as quite broad ragged doublet around 2.0 ppm. In all complexes, peaks in therange 1.20–1.60 ppm are from co-ordinated water. The various assignments of 1H NMRof the ligand and metal (II) complexes are summarized in Table 3. Chemical shift are inppm from TMS & multiplicity in parentheses (bd, broad; d, doublet; m, multiplet).

3.3. Electronic spectra and magnetic susceptibility

The electronic spectra of the ligands and its metal complexes have been studied in therange 190–800 nm. The shoulder band observed at 275 nm in water solvent in ligandsmay be assigned to n ! n* transition within the OH group of the hydroxyl moiety in thefree ligands. This band disappeared in all the complexes, revealing the involvement of theOH of hydroxyl group oxygen in chelate formation. The band observed at 296 nm in GM(Gentamicin) may be assigned to p ! p* transition of the oxygen of hydroxyl groupof ligand. In the cobalt complex, the absorption band in the visible region 508–565 nm and470–430 nm, assignable to 2B2g ! 4Eg(P) and

2B2g ! 4A2g(P) transition, suggests squareplanar geometry of the cobalt complex, whereas in the nickel complex, absorption bands

Table 3. 1H NMR assignments of ligand and its Ni(II) complex.

Assignment Ligand (2)

NH 2.0 (d) 8.19 (bd)C1H, C2H 3.26 3.26 (m)C(20)H, C(60)H 1.94, 2.03, 1.92 1.94, 2.03, 1.92 (d)C(1)OH 4.81 –C(3)OH 4.81 (bd) –C(5)H 3.26 6.28C(1)H 2.03, 1.92 5.03C(2)H 1.94 (m) 4.40 (m)C(3)H 3.70 (m) 3.68 (m)CH(H) 3.12, 2.73, 4.73 4.95, 4.73 (m)

Table 4. Thermodynamic activation parameters of the metal complexes.

ComplexOrder/n Steps

E*/Jmol71 A/s71

DS*/JK71

mol71DH*/

J mol71DG*/

kJ mol71

[Co(L)H2O] (1) 1 I 45.34 5.25 6 105 711.56 41.14 53.93II 39.43 2.81 6 105 756.67 34.87 35.41

[Ni(L)H2O] (2) 1 I 34.09 8.27 6 105 725.096 84.10 96.62II 39.75 2.16 6 105 7146.69 1205.89 106.53

[Cd(L)H2O] (3) 1 I 577.78 6.99 6 104 7159.46 22.95 52.97II 34.11 4.08 6 104 7166.13 59.79 98.87

[Sn(L)H2O] (4) 1 I 59.18 3.28 6 106 7123.65 34.01 39.97II 25.82 2.42 6 106 7133.59 1340.06 90.31

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in the visible region 723–657 nm and 580–560 nm may be assigned to the 1A1g !1A2g and

1A1g !1B1g transitions, respectively, which are consistent with a square planar

stereochemistry with the nickel (II) ions [22,23]. The magnetic moment of the cobaltcomplex is 2.7 BM at room temperature. The reason for the departure from the spin onlyvalue lies partly in the existence of the second order Zeeman Effect between the groundand higher ligand field terms [24]. However, it lies mainly in fact that in the presence ofspin orbit coupling the quenching effect of the ligand field cannot be complete. Nickel,cadmium and lead complexes are diamagnetic in nature.

3.4. Kinetics of thermal decomposition

Recently, there has been increasing interest in determining the rate-dependent parametersof solid-state non-isothermal decomposition reactions by analysis of TG curves [25,26].

Figure 2. TOF Mass-Spectra of (2).

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Thermogravimetric (TG) and differential thermo gravimetric (DTA) analyses were carriedout for different metal–tetracycline complexes in ambient conditions. The thermogravi-metric analysis revealed that the complexes of cobalt and nickel lose mass between 658Cand 1408C, corresponding to nearly 13% of the total mass, followed by considerabledecomposition up to 6008C, which corresponds to the decomposition of the ligandmolecule leaving metal oxide (CoO and NiO, respectively) as residue. The complexes ofcadmium and tin decomposes nearly 9% of the total mass up to temperature 1708C,

Figure 3. TOF Mass-Spectra of (4).

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followed by considerable decomposition of the ligand molecule up to 6508C, leaving metaloxide (CdO and SnO2, respectively) as residue. On the basis of thermal decomposition, thekinetic analysis parameters such as activation energy (E*), enthalpy of activation (DH*),entropy of activation (DS*), free energy change of decomposition (DG*) were evaluatedgraphically by employing the Coats–Redfern relation [27]

Log ½�Log ð1�aÞ=T2 ¼ log ½AR=y E�ð1� 2RT=E�Þ� � E�� 2:303RT ð1Þ

where a is the mass loss up to the temperature T, R is the gas constant, E*is the activationenergy in J mol71, y is the linear heating rate and the term (1 7 2RT/E*) ffi 1. A straightline plot of left hand side of the Equation (1) against 1/T gives the value of E* whereas itsintercept corresponds to A (Arrhenius constant). The Coats and Redfern linearizationplots, confirms the first order kinetics for the decomposition process [28]. The calculatedvalues of thermodynamic activation parameters for the decomposition steps of the metalcomplexes are reported in Table 4. According to the kinetic data obtained from the TGcurves, the activation energy relates the thermal stability of the metal complexes. Amongmetal complexes, activation energy increases as complex III * complex II 5 complexIV 5 complex I, same trends happens with thermal stability of metal complexes. All thecomplexes have negative entropy, which indicates that the complexes are formed sponta-neously. The negative value of entropy also indicates a more ordered activated state that maybe possible through the chemisorptions of oxygen and other decomposition products. Thenegative values of the entropies of activation are compensated by the values of the enthalpiesof activation, leading to almost the same values for the free energy of activation.

Table 5. Crystallographic data for the metal complexes.

Empirical formula [Co(L)H2O] (1) [Ni(L)H2O] (2) [Cd(L)H2O] (3) [Sn(L)H2O] (4)

Formula weight 570.526 570.226 625.23 630.306Temperature (K) 298 298 298 298Wave length (A) 1.54056 1.54056 1.54056 1.54056Crystal system Orthorhombic Monoclinic Monoclinic TriclinicSpace group P nma P 4/m P 2/m P1Unit cell

dimensions (A)a ¼ 14.045140,b ¼ 9.706635,c ¼ 5.50331,a ¼ 908,

b ¼ 90.008,g ¼ 908

a ¼ 16.679200,b ¼ 12.714760,c ¼ 5.267300,

a ¼ 90,b ¼ 95.21,g ¼ 908

a ¼ 11.129740,b ¼ 10.460970,c ¼ 8.124669,

a ¼ 90,b ¼ 104.6368,

g ¼ 908

a ¼ 8.295049,b ¼ 9.619381,c ¼ 16.12853,a ¼ 63.75208,b ¼ 53.3017,g ¼ 62.28692

Volume (A3) 750.55 1112.00 915.24 888.682y Range 10–708 5–608 10–708 10–708

Limiting indices 0 � h �8,0 � k �4,0 � l �4

77 � h �7,0 � k �6,0 � l �2

75 � h �4,0 � k

�5, 0 � l �4

72 � h �3, 73 � k �4,0 � l �5

Density (g/cc) 1.144 1.151 1.073 1.48Z 1 2 1 2

Avs. Eps 0.0000676 0.0000687 0.0000692 0.0000424Merit of fitness M (12) ¼ 7,

F (12) ¼ 3M (20) ¼ 7,F (20) ¼ 5

M (20) ¼ 7,F (20) ¼ 6

M (15) ¼ 6,F (15) ¼ 3

Reflections 667 1123 1224 576Particle size (nm) 33.812 9.97.74 8.591 2.79

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3.5. TOF–MS spectra

Mass spectrometry has been successfully used to investigate molecular species [MH]þ insolution [9–11]. The molecular ion peaks of the ligands and complexes have been used toconfirm the proposed formula (Table 4). The pattern of the mass spectrum gives animpression of the successive degradation of the target compound with the series of peakscorresponding to the various fragments. Their intensity gives an idea of stability offragments. The ligand starts degradation and finally forms [C13H24N3O6]e,317/318 (100%m/z values). In the TOF–mass spectra of metal complexes initial fragmentation pattern isagain similar (loss of two water molecules), a mononuclear nature for these complexes[M(L)]þ can be deduced. The last two fragments appears in nearly all the complexes at

Figure 4. (a) Optimized structure of complex (2); (b) Stereo structure of complex (2).

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positions (m/z values) 165(100%) (1), 100% (2), 100% (3) and 57% complex (4) and 223/225/227 [10% (1), 50% (2) and 100% (4)] corresponds to [C8H7NO3]

þ and [C9H8NO6]þ

respectively, which could be the result of degradation and demetallation of the complexes(Figures 2 and 3). Scanned TOF-MS spectrum of (2) and (4) with specific fragments.

3.6. X-ray powder diffraction studies

In the absence of single crystal, X-ray powder data are especially useful to deduce accuratecell parameters. The diffraction pattern reveals the crystalline nature of the complex. The

Figure 5. (a) Optimized structure of complex (4); (b) Stereo structure of complex (4).

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indexing procedures were performed using (CCP4, UK) Crysfire programme [29] givingdifferent crystal system with varying space group. The merit of fitness and particle size ofthe metal (II) complexes has been calculated. The crystallographic data of the complexesare shown in Table 5.

Figure 6. Graphical structure of (4).

Table 6. Data for selected bond lengths and bond angles of complex 2 and complex 4.

Bond length (A8) Bond angle (8)

Complex (2) Complex (2)O(25)-Ni(21) 1.7804 H(59)-O(23)-Ni(21) 99.0995O(24)-Ni(21) 1.7796 H(58)-O(23)-Ni(21) 99.6118Ni(21)-O(23) 1.8131 H(57)-O(22)-H(56) 71.5181Ni(21)-O(22) 1.8479 H(57)-O(22)-Ni(21) 93.0445Complex (4) H(56)-O(22)-Ni(21) 164.3067O(25)-Sn(21) 2.0520 O(25)-Ni(21)-O(24) 115.1903O(24)-Sn(21) 2.0535 O(25)-Ni(21)-O(23) 114.7300Sn(21)-O(23) 2.0824 O(25)-Ni(21)-O(22) 77.8340Sn(21)-O(22) 2.0743 O(24)-Ni(21)-O(23) 115.5157

O(24)-Ni(21)-O(22) 81.7984O(23)-Ni(21)-O(22) 71.6623

Complex (4)H(59)-O(23)-Sn(21) 170.0491H(58)-O(23)-Sn(21) 94.6840H(57)-O(22)-Sn(21) 96.1848H(56)-O(22)-Sn(21) 96.6765O(25)-Sn(21)-O(24) 113.1224O(25)-Sn(21)-O(23) 107.9120O(25)-Sn(21)-O(22) 109.7941O(24)-Sn(21)-O(23) 106.0005O(24)-Sn(21)-O(22) 110.4817O(23)-Sn(21)-O(22) 109.3812

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3.7. Molecular structures and analysis of bonding modes

In order to ascertain the structural preferences and coordination behavior of gentamicin tometal ions, molecular mechanics calculations on the [ML]nþ species were undertaken. Wehave optimized the molecular structure of complexes (Figures 4 and 5). The optimizedmolecular structure of (2) is given in Figure 4(a). Energy minimization was repeated severaltimes to find the global minimum. The energy minimization values for the optimizedstructure for (2) and (4) are 35.12 and 25.31 kcal/mol, respectively. The selected bond lengthand bond angles of (2) and (4) are represented in Table 6, is comparable with square planargeometry. This supports square planar geometry of the metal complexes (Figure 6).

4. Conclusion

It can be concluded that the metal(II) complexes of gentamicin are monomeric and involvecoordination through the tertiary alcoholic OH group, giving square planar geometries.Spectroscopic data in solution showed that the gentamicin adopts similar coordinationbehavior, as suggested by vibrational spectra. Cell parameters and optimization ofgeometry of the metal complexes were also carried out.

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