- Email: [email protected]
TD-DFT theoretical studies, experimental characterization, electrochemical and antioxidant activity of Fe(III) complexes of bis (dimethylglyoximato) guanine
Journal of Molecular Structure 1186 (2019) 413e422
Contents lists available at ScienceDirect
Journal of Molecular Structure journal homepage: http://www.elsevier.com/locate/molstruc
Synthesis, DFT/TD-DFT theoretical studies, experimental characterization, electrochemical and antioxidant activity of Fe(III) complexes of bis (dimethylglyoximato) guanine Lamia Abane-Merzouk a, Ahmed Adkhis a, *, Souhila Terrachet-Bouaziz b, Malika Makhloufi-Chebli a a b
Laboratoire LPCM, D epartement de Chimie, Facult e des Sciences, Universit e Mouloud Mammeri, 15000, TiziOuzou, Algeria Department of Chemistry, Faculty of Sciences, University Mohamed Bouguerra, Boumerdes, Algeria
a r t i c l e i n f o
a b s t r a c t
Article history: Received 13 January 2019 Received in revised form 21 February 2019 Accepted 26 February 2019 Available online 4 March 2019
A new series of Iron (III) complexes of general formula [Fe(Hdmg)2(Gua)X]nH2O where (Hdmg ¼ dimethylglyoximato monoanion, Gua ¼ Guanine, X ¼ Cl, Br or I) have been prepared and characterized by elemental analysis, conductivity, infrared spectra and electronic spectra. The molar conductance data confirm that all the complexes are no electrolytic. The mode of bonding and geometry of the Fe(III) complexes have been inferred spectral data, an octahedral geometry is ascribed for all the Fe(III) complexes. In these compounds, two dimethylglyoximato monoanions bond to the metal in the equatorial plan have been formed. The two axial sites being occupied by the guanine and the halogen moieties. The thermal properties of the Fe(III) complexes were investigated using TG and TDA. The thermal behavior of the compounds illustrates the general decomposition patterns of the complexes. The voltammetry experiments have been performed on the complexes to evaluate the redox process of Fe(III)/Fe(II) couple. These compounds were screened for their in-vitro antioxidant properties using 2,2diphenyl-1-picrylhydrazyl radical (DPPH .) free radical scavenging. The results obtained show that these complexes have a good antioxidant activity in comparison with ascorbic acid as positive control. © 2019 Published by Elsevier B.V.
Keywords: Dimethylglyoxime Iron(III) complexes Spectroscopic analysis Cyclic voltammetry Anti-oxidant activity Drug-likeness ADME analysis
1. Introduction Iron is the most abundant metal present at active sites of a large number of metalloenzymes [1]. The synthesis and characterization of iron complexes with different ligands play an important role in the coordination chemistry of iron because of their magnetic properties which exhibit spin states, low-spin and high-spin [2]. Furthermore it was found that some iron(III) complexes provide a useful structural and electronic model for the similarly coordinated iron(III) sites found in the hem iron enzymes [3,4]. On the other hand, oximes are well known for their high coordinating abilities and versatile behavior, but much remains to be learned about the type of structures that are formed as well as the factors that dictate structure. In general, dioximes coordinate with metal ions in four ways [5]. The most stable form corresponds to that shown in Fig. 1. This stability is due to the existence of the hydrogen bonds which
* Corresponding author. E-mail address: [email protected] (A. Adkhis). https://doi.org/10.1016/j.molstruc.2019.02.108 0022-2860/© 2019 Published by Elsevier B.V.
are formed between the hydroxyl groups. The dimethylglyoxime ligand (dmgH2) is a particulary interesting chelating oxime ligand [6]. The general characteristic of the structure of transition metal dioxime complexes are summarized early in the works of Chakravorty [7]. It is well known that the nucleotides acids vary in their ability to participate in many biological reactions forming metal complexes [8]. The diversity of research contribution has been affected by a large variety of studies concerning the chemical and biological aspects of the interactions between metal ion and nucleic acids and their constituent [9]. In recent years, interest in natural or synthetic antioxidants, in relation to therapeutic properties, has increased considerably. Antioxidants are the subject of numerous researches due to their importance in the health of pernicious diseases (cancer) and in the agri-alimo industry. Compounds containing pyrimidine and purine play a significant role in many biological systems, where both exist in nucleic acids, several vitamins, coenzymes and antibiotics [10]. These provide potential binding sites for metal ions, and any information on their coordinating properties is important as a means of understanding the role of the metal ions in biological systems.
414
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422
ethanolic solution (40 ml) and H2dmg (2.32 g, 20 mmol) dissolved in a hot ethanol (60 ml) was added powder guanine (1.51 g, 10 mmol). The reaction mixture is stirred and refluxed for approximately 4 h at a temperature of 50 C. The brown precipitated solid obtained was collected by filtration, washed with ethanol and diethylether. The solid was dried in air. Yield: 45%; Anal. Calc. for C13H33FeN9O11Cl: Fe, 7.43%; C, 27.02%; N, 21.82%; H, 7.27%. Found: Fe, 7.73%; C, 27.23%; N, 21.85%; H, 7.24%.
Fig. 1. Structure of dimethylglyoximato complexes M(Hdmg)2.
Metal ion-nucleic acid complexes and metal oxime complexes have extensive applications in the fields of medicine, photo and magnetic chemistry and electrochemistry. In addition, iron is needed by the bacteria and different iron complexes of bioactive ligands have shown anti-microbacterial activity [11]. Hence in this work we report the syntheses, characterization of iron(III) complexes, and their antioxidant activities.
2.2.2. Synthesis of [Fe(Hdmg)2(Gua) Br]4H2O (complex 2) To an aqueous solution (40 ml) of Fe(NO3)3.9H2O (4.04 g, 10 mmol) and NaBr (1.02 g, 10 mmol) was added dropwise a hot ethanolic solution (60 ml) of dimethylglyoxime (2.32 g, 20 mml). Then powdery guanine (1.51 g, 10 mmol) was added to the mixture. The resulting mixture is left stirring under reflux at 50 C for 3 h. The brown precipitate obtained was isolated by filtration and washed with ethanol and diethylether. The solid was dried at room temperature. Yield: 37%; Anal. Calc. for C13H29FeN9O9Br: Fe, 8.39%; C, 26.63%; N, 21.5%; H, 4.09%. Found: Fe, 8.47%; C, 26.57%; N, 21.20%; H, 4.02%.
2. Experimental
2.2.3. Synthesis of [Fe(Hdmg)2(Gua)I]3H2O (complex 3) This compound was synthesized using the same procedure as described above for complex II, starting from KI (1.66 g, 10 mmol) instead of NaBr. Yield: 42%; Anal. Calc. for C13H27FeN9O8I: Fe, 9.36%; C, 25.38%; N, 20.49%; H, 3.57%. Found: Fe, 9.78%; C, 25.94%; N, 20.04%; H, 3.37%.
2.1. Materials and methods
2.3. Antioxidant activity
The chemicals reagents and solvents (Fluka products) used were of analytical grade and were used without further purification. Melting points were determined on a Stuart scientific SPM3 apparatus fitted with a microscope and are uncorrected. The infrared spectra were recorded in the region 4000e400 cm1 on a BRUKER TENSOR 27 IR spectrophotometer. Electronic spectra were measured on a thermo Scientific EVOLUTION 220 ultravioletevisible spectrophotometer (in DMSO solution); measurements were made from 200 to 800 nm. The elemental microanalysis (C, H, N) was carried out at the Service of analysis and characterization of CNRS (Poitier France. The molar Conductance values were obtained for 10 3M in DMSO solution at 25 ± 1 C determined using melting point meter MPM-H2. All electrochemical experiments were carried out using a potentiostatgalvanostat EGG-273 A controlled with Power-suite software. The electrochemical cell was equipped with a modified carbon past disk as the working electrode, a platinum electrode as counter electrode and silver/silver chloride (Ag/AgCl) electrode as the reference electrode. The electrolytic bath consisted of the complexes in DMSO solution (103 M) and sodium perchlorate (NaClO4 101 M) was used as a supporting electrolyte to increase the conductivity of the bath. The electrochemical characterization was performed using the cyclic voltammetry (CV) with a scan rate of 50 mV/s. TG/DTA/ EDS techniques were obtained with NETZSCH STA 409 PC/PG. The heating rate of 10 C/mn was used between 25 C and 600 C. This theory uses Becke's hybrid three-parameter method and the Lee-Yang-Parr correlation function (B3LYP) associated with the use of ORCA V4 software, which has proved very effective for a large number of organic and inorganic systems.
The antioxidant activity (free radical scavenging activity) of the synthesized H2dmg metal complexes was evaluated using the 2,2diphenyl-1-picrylhydrazyl free radical (DPPH.) scavenging assay [12,13]. DPPH. is a stable free radical which has been widely accepted as a tool for estimating the free radical scavenging activities of potential antioxidants (Fig. 2) [14]. The free radical scavenging ability of H2dmg metal complexes was studied spectrophotometrically using DPPH radical scavenging technique. Radical scavenging activities of all complexes were determined by the interacting ability of compounds with stable free radical DPPH. The degree of discoloration indicated the scavenging potential of the antioxidant compounds in terms hydrogen donating ability. DPPH. solution was prepared by dissolving DPPH. in ethanol to give a concentration of 4 mg/100 mL1. The metal complexes were dissolved in DMSO to obtain a solution of 101 M. The concentration of test compounds was diluted with DMSO to get final concentrations 0.05, 0.025 et 0.0125 mol/l for all the compounds, but the standard were further diluted to give concentration solutions of 0.1, 0.05, 0.025 et 0.0125, 0.00625, 0.003125, 0.0015625 mol/l. 40 mL of each tested concentration of each complex was added to each well separately in diplicate, and then 2 mL of DPPH solution was added. Negative control wells were loaded with 40 mL of DMSO and 2 mL of DPPH. solution each. After vortexing, the mixtures were incubated at room temperature for 1 h in darkness at 25 C, and
2.2. Synthesis of iron(III) complexes 2.2.1. Synthesis of [Fe(Hdmg)2(Gua)Cl]6H2O (complex 1) To a mixture of FeCl3.6H2O (2.70 g, 10 mmol) dissolved in aqua
Fig. 2. Structure of the DPPH. radical and its DPPH2 reduction product.
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422
then the absorbance of these compounds at different concentrations was recorded at 517 nm. The reduction of the DPPH radical was measured by monitoring continuously the decrease of absorption at 517 nm. DPPH scavenging effect was calculated as percentage of DPPH discoloration using the equation:
RSA ð%Þ ¼ ½ðAc AsÞ=Ac x 100 where Ac is the absorbance of the control (absorbance of DPPH. ethanol solution without sample), and As is the absorbance of the tested compound after 60 min incubation. Minimum inhibitory concentrations IC50 of the tested compounds are determined from RSA (%) against concentration graph. 3. Results and discussion The complexes were obtained as powders and attempts to obtain a single crystal suitable for X-ray determination were unsuccessful. Their physical properties and analytical data are summarized in Table 1 and their empirical formulas were obtained based on elemental analysis and TGA measurements. The complexes are insoluble in water and soluble in DMF and DMSO. The results of elemental analysis suggested that the complexes can be formulated as [Fe(Hdmg)2(Gua)X]nH2O, showing that dimethylglyoximato group as a monoanion. The melting point of all complexes exceeds 360 C indicating that all complexes are stable in air and can be preserved at room temperature. The molar conductance values of all complexes in DMSO at 103 M supporting the non-electrolyte nature of the all iron complexes. 3.1. UVevisible spectroscopic study on the effect of the stoichiometric amount of the metal on the ligand To optimize the complexation, a study was carried out to determine the molar ratio M/Ls required for the synthesis of the complexes. The solution of the ligands and that of the metal were prepared in Ethanol-water (V/V) at concentrations of 105 M for the ligands and 0.125 105, 0.25 105, 0.5 105, 0.75 105 and 1 105 M for the Fe(NO3)3.9H2O. The absorption spectrums of free ligands H2dmg and guanine and of the mixture H2dmg-guanine (2:1) have been studied (Fig. 3AeC). In the case of the mixture H2dmg H2dmg-guanine, two bands were observed and the intense long-wavelength band situated around 231 nm was attributed to the p - p* transition. In the O present work, addition of Fe(NO)3.9H 3 2 to a solution of H2dmg-guanine 5 (2:1) (10 M) in ethanol-H2O (V/V) induced drastic changes in the absorption spectrum. A strong hypochromic effect on the wavelength band at 231 nm (Fig. 4, curves a to f) was observed. The band around 302 nm underwent a hypechromic effect. The absorbance was analysed as a function of cation concentration (Fig. 5). Since the shape of the absorption spectrum varied strongly in the presence of cations, the absorbance was recorded at two different wavelengths chosen in order to obtain maximum
415
information. The calculation was performed simultaneously on the two bands obtained by absorption spectroscopy. Good fits were obtained, as displayed in Fig. 5. The results of Fig. 4 show that from an equimolar ratio metal/ guanine and 2 equivalents of H2dmg, the absorbance becomes important. The time factor on metal (1eq)-ligands (H2dmg-guanine (2:1 eq)) coordination (1: 2:1) is also studied. Hypechromic effect on the wavelength bands at 231 nm and 302 nm is observed (Fig. 6). 3.2. Characterization of the ligands and the complexes 3.2.1. Infrared spectra In the absence of X-ray structure determination, infrared spectra are the most suitable technique to determine the coordinating atoms of the ligands by comparing the free ligand spectra with those of the complex. The main vibrations of the ligands and the complexes are presented in Table 2. The IR spectra for complexes 1, 2 and 3 are rather similar; the discussion is confined to the most important vibrations. The IR spectra of the complexes were compared with the spectra of H2dmg and guanine to establish the coordination modes of the ligands in the complexes. The spectra of all the complexes show bands in the region 3335-3338 cm1 due to the NH2 stretching vibration of guanine which is shifted to higher frequency that encountered in the free guanine (3322 cm1), the spectrum of the chloro complex is given by way of example in Fig. 7. This considerable shifts in ʋ(NH2) clearly indicate the involvement of NH2 of guanine in metal coordination [15]. The guanine exhibit a ʋ(CH) stretching vibration bands at 2991 cm1 and 2900 cm1. In the spectra of the complexes, these bands are located at around 2995 cm1 and 2902 cm1. The bands are apparently insensitive to the complexation [16]. Characteristic IR strong bands of free guanine corresponding to ʋ(C¼O) observed at 1693 and 1670 cm1, are shown in spectra of the complexes at 1690e1699 cm1 and 16611667 cm1 region. The little negative shifts indicate that the oxygen atom of carbonyl function is not participating to the metal coordination [17]. An additional band which appears at 3445, 3450 and 3465 cm1 respectively in the spectra of all the complexes is correlated with the ʋ(OH) vibration of water hydratation molecules. In the infrared spectra of H2dmg, the sharp band of medium intensity occurring at 1447 cm1 is attributed to y (C¼N) of the oxime. In the complexes, this band is shifted to higher frequency (1557 cm1) [18]. This suggests that dimethylglyoxime is coordinated to the metal ion through the nitrogen atom of the oxime. In all the complexes, the strong as. Well as sharp band between 1208 and 1206 cm1 is due to the NeO stretching vibration. The metal sensitive as sharp band at 772779 cm1 region may be attributed to the C¼NeO deformation vibration [19]. The peak around 1363 cm1 in the spectra of H2dmg is assigned to the deformation of -CH3 groups; this peak is observed in 1378e1387 cm1 region in the spectra of the compounds [20]. The characteristic stretching vibrations belonging to y(OH) of
Table 1 Analytical data of the complexes. M.C (U1.cm2. mol1)
Compound Formula
Color
Mol.Wt (g/mol)
Yield (%)
M.P (C )
Anal-Found/Cal (%) C
H
N
Fe
Fe[(Hdmg)2(Gua)Cl]6H2O
Brown
577.66
45
>360
Fe[(Hdmg)2Gua)Br]4H2O
Brown- Russ
586.11
37
>360
[F Fe[(Hdmg)2 (Gua)I]3H2O
Brown
615.11
42
>360
27.23 27.02 26.57 26.63 25.94 25.38
7.24 7.27 4.02 4.09 3.37 3.57
21.85 21.82 21.2 21.5 20.04 20.49
7.73 7.43 8.47 8.39 9.78 9.36
4.14 4.28 5.27
416
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422
Fig. 4. Absorption spectrum of the ligands (105 M) in ethanol in the absence (curve a) O Fe(NO) .9HO and presence of Fe(NO)3.9H 3 2 . Curves b to f: effect 33 2 addition. From bottom to top: Fe(NO) 9HO 5 5 5 5 and 2.105 M. 33 2 : 0.125.10 ; 0.25.10 ; 0.5.10 ; 10
Fig. 5. Fitted spectrophotometric data for the ligands vs. cation concentration. Variation of absorbance: l ¼ 231 and 304 nm.
the complexes 1, 2 and 3 is the absorption at 509e506 and 503 cm1 respectively; this was assigned to the (FeeN) vibration. These bands are not found in the spectrum of the ligand [22]. The variation in the position of these bands follows the order in spectrochemical series: I < Br < Cl. The representative IR spectra of [Fe(Hdmg)2(Gua)Cl]4H2O complex is shown in Figure and the other spectra are given in the Supplementary material section.
Fig. 3. (A) Absorption spectrum of H2dmg (105 M); (B) Absorption spectrum of guanine (105 M); (C) Absorption spectrum of H2dmg eguanine(2:1) (105 M).
oxime is observed at 3205 cm1 in dimethylglyoxime, this band is expected to disappear upon complex formation. The band in the region 947-951 cm1 may be attributed to deformation vibration of OH in H2dmg moiety [21]. The only new feature in the IR spectra of
3.2.2. Electronic spectra The electronic absorption spectrum is often very useful on the evaluation of structural investigation. The electronic absorption spectra of Fe(III) complexes was recorded at room temperature using DMSO as solvent. The spectra of H2dmg, guanine and its complexes with tentative assignments are presented in Table 3. The H2dmg exhibit two absorption bands in the UV, the first band with a strong intensity appear at 227 nm, the second lower intensity band appear at 313 nm (Fig. 3A). These bands can be attributed to p / p* transition of the mimic (C¼N) function present
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422
417
Fig. 7. Experimental FT-IR spectra of [Fe(Hdmg)2(Gua)Cl]4H2O. Fig. 6. Absorption spectrum of the ligands (105 M) in ethanol in the presence of Fe(NO) 9HO 5 M) along the time. 33 2 (10
in the ligand moiety. The spectra of guanine display in the UV an intense band at 273 nm which could be ascribbed to p / p* transition (Fig. 3B). The electronic spectra of the three complexes show intense bands at 289, 304 and 296 nm respectively, which could be assigned to p / p* transition. The position of these bands is shifted to higher wavelength (lower energy) compared to that observed in the free ligand (for example UVevisible spectra of complex 1, Fig. 8). The chloro, bromo and iodo dimethylglyoximato Fe(III) complexes exhibit similar spectra in the UV and in the visible region. While the spectral features observed in the range 322e376 nm may be assigned to X Fe(III) (X ¼ Cl, Br or I) charge transfer transition [23]. These values show that the general trend observed is shifting of the L/M transition bands in the complexes increases with the growth of electronegativity of the axial ligand: I < Br < Cl [24]. In addition, the very relatively low intense band appear at around 473 nm assignable to the spin allowed electronic transition corresponding to 6A1g / 4T1g (G) transition consistent with octahedral Fe(III) complexes [25,26]. The results obtained describe a pseudo-octahedral structure for all iron complexes (Fig. 9), where the two dimethylglyoxime monoanions are in the equatorial plane. The guanine and X (halogen) occupy the two axial coordination sites.
Table 3 Electronic spectra data of iron (III) Complexes. Complexes
l(nm)
ῡ(cm1)
Electronic transition
H2dmg Guanine
227 313 273 289 376 486 304 328 477 296 322 356 473
44053 31949 33630 34602 26596 20576 32894 30487 20964 33784 31056 28090 21141
p - p* p - p* p - p* p - p*
Complex 1
Complex 2
Complex 3
C.T. Cl/Fe(III) Transition d-d p - p* C.T. Br/Fe(III) Transition d-d p - p* C.T. I /Fe(III) Transition d-d
3.3. Complex structures and theoretical studies In order to determine theoretically the geometry of the complexes and make a comparison with the experimental observations, quantum chemical optimisations of the I. R and UVeVISIBLE have been carried out. The geometries of the synthesized complexes were optimized with density functional theory (DFT) using Becke's three parameters hybrid method and the Lee-Yang-Parr correlation functional (B3LYP) [27e30] combined with LANL2DZ [31e34] using ORCA v4
Table 2 Relevant infrared spectra bands for the ligands and Complexes (cm1). Complexes
y(OH)
H2dmg Guanine
3205
Complex 1
e
Complex 2
e
Complex 3
e
y
as(NH2) s
y (NH2)
3320 3116 3335 3115 3335 3109 3338 3112
n (CH)
y(C¼O)
n (CN)
n(NO)
d (NO)
n (FeeN)
e 2991 2905 2993 2904 e 2903 2995 2902
e 1693 1670 1694 1669 1699 1661 1690 1667
1447 e e 1557
1144 e e 1208
750 e e 776
e e e 509
1557
1206
772
506
1557
1208
779
503
418
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422
Fig. 10. Theoretical electronic Spectra of complex 1 (X ¼ Cl).
Fig. 8. UVevisible spectra of the complex [Fe(Hdmg)2(Gua)Cl]6H2O in DMSO.
Fig. 9. Suggested structure of the complexes. (X ¼ Cl, Br or I). Fig. 11. Theoretical FT-IR spectra of complex 1 (X ¼ Cl).
software [35,36]. Infrared spectra were obtained at the same level of theory combined with Time-Dependent Density Functional Theory (TD-DFT) for UVevis absorption. The calculated UVevis absorption is in agreement with the experimental results. It is worth noting that previous studies were demonstrated the capability of chosen method and basis set to predict the electronic properties of metallic complexes quite reasonably [37e41]. For example, the calculated absorption spectrum of the complex [Fe(Hdmg)2(Gua)Cl]4H2O (Fig. 10) shows in the ultraviolet domain a band around 320 nm due to the transition p / p *. Another band located at 430 nm corresponds to the metalligand charge transfer. A band appears in the visible which correspond to the transition 6A1g / 4T1g. The ligands coordination sites which are involved in bonding with the metal ions have been determined by careful comparison of the infrared absorption spectra of the complexes with that of the parent ligand. The vibrational assignments were carried out with support DFT calculations. The theoretical FT-IR spectra of [Fe(Hdmg)2(Gua)Cl]4H2O is given in Fig. 11 and compared with the experimental spectra. The IR spectra indicate that the ligands are coordinated with the metal, resulting in neutral species. Coordination occurs through the N-atoms from the H2dmg and Guanine moiety and from the
halogen. On basis of the elemental analysis, electronic spectra and theoretical studies, the complexes were formulated as Fe [(Hdmg)2(Gua)Cl]6H2O, Fe[(Hdmg)2Gua)Br]4H2O and Fe[(Hdmg)2 (Gua)I]3H2O (Fig. 12). 3.4. Electrochemical: cyclic voltammetry study The application of electrochemical methods to the study of the coordination of metal ion provides a useful complement to the previously used methods of investigation, such as UVevisible spectroscopy. The electrochemical behaviours of the ligands and complexes were examined by cyclic voltammetry. DMSO solution (103 M) containing 101 M sodium perchlorate monohydrate (NaClO4) as supporting electrolyte was used. The potential were scanned in the range [-1,4 V, þ1,4 V] at a scan rate of 50 mV s1. The cyclic voltammogram of guanine display two cathodic waves at 0.78 V and 0.47 V of the molecule, one anodic waves is observed at þ0.75 V around, due probably to the oxidation of guanine. In the voltammogram of H2dmg, one anodic wave is located at þ0.95 V
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422
419
Table 4 Cyclic voltametry data for the iron(III) complexes. Complexes Complexes
Couple red/ox
Epc[V]
Epa[V]
DEP[mV]
ipa/ipc
Complex 1
Fe(III/II) Fe(III/II)
Complex 3
Fe(III/II)
0.52 þ0.26 0.58 þ0.23 0.57 þ0.57
500 370 120 530 150 e
0.38
Complex 2
1.02 0.11 0.70 0.30 0.72 e
0.27 0.42 0.10 e
assigned to the oxidation of the oxime group [42]. The reduction potential and the corresponding values summarized in Table 4, correspond to one electron irreversible metal centered Fe(III)/Fe(II) redox processes. The representative cyclic voltammogram of [Fe(Hdmg)2(Gua)Br] 4H20 complex is shown in Fig. 13. The complexes exhibited irreversible one electron response (DE ˃ 100 m V) corresponding to the couple Fe(III)/(II) [43]. This observation can be explained by the break of the bonding Fe (III)-X lasting the reduction of the complexes [Fe(III)(Hdmg)2GuaX]. So the axial ligand (X) is lost, and the formation of pentacoordined complex of Fe(II) according to the reaction (*) can be envisaged without the axial ligand X. [Fe / [Fe(II)(Hdmg)2 (Gua)] þ X (X ¼ Cl, Br, (III)(Hdmg)2(Gua)X] þ e I).(*) [44]. Values in Table 4 show that for all the complexes, the current ratio of the direct and reverse reduction waves is less then unity. This indicates the low chemical stabilities of this species in time scale. The heigh stability of the reduced forms of the complexes is also confirmed by the high DEp value which is compelling evidence for the irreversibility of electron transfer reaction. The cyclic voltammetry of complexes I and II presents also two defined cathodic waves at 0.11 V and 0.30 V, respectively and two anodic waves at þ0.26 V and þ0.23 V. The redox processes must be considered irreversible. In the other hand, defined oxidation processes in the cyclic voltammetry of the complexes II and III is observed at positive potentials þ0.56 and þ 0.57 V respectively, these potentials can be attributed to oxime function in the complexes. 3.5. Thermal study The experimental data corresponding to the chloro, bromo and iodo complexes are listed in Table 5. The thermograms indicate that all complexes involved water molecules [45]. The three complexes show decomposition pattern of three stages.
Fig. 12. Optimized Geometry for complexes 1 (A); 2 (B) and 3 (C). Fig. 13. Cyclic voltammetry curve of [Fe(Hdmg)2(Gua)Br]4H2O.
420
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422
Table 5 Thermal decomposition data of the Fe (III) Complexes. Compound
Temperature Range (C)
Weight loss (Obs)
Weight loss (Cal)
Species formed
Complex 1
25e120 120e300 300e600 25e160 160e380 380e600 25e120 120e320 320e600
3.13 20.94 62.16 6.44 22.79 38.32 5.74 32.62 40.00
3.11 19.24 62.59 6.14 21.06 38.32 5.85 32.68 41.60
[Fe(Hdmg)2(Gua)Cl]5H2O [Fe(Hdmg)2(Gua)Cl0.56] Fe(Hdmg)1.75
Complex 2
Complex 3
For the chloro complex, the first steps with estimated mass loss of 3.13%, found within the temperature range 25e120 C, corresponding to loss of uncoordinated water molecule. He second steps with estimated weight loss of 20.94% within the temperature rang 120e300 C, can be attributed to the liberation of five latter water molecules and approximately moiety of the chloride in the structure [46,47]. The third step is started at 300 C and ends at 600 C, the corresponding mass loss (62.12%) is due the decomposition of the guanine molecule and a fragment of Hdmg. The mass loss in the third stage is in agreement with the masse calculated loss of 62.59%. The bromo complex exhibit one weight loss step over the temperature ranges 25e120 C, due to the elimination of the two water molecules of hydration. The second step which corresponds to an endothermic peak corresponds to the elimination of two water molecules hydration and the bromide. The third thermal reaction is endothermic within 380e600 C range. The mass loss in this stage corresponds approximately to the guanine and 0.21 Hdmg. In iodo complex, the single step weight loss corresponds to simultaneous disappearance of two water molecules hydration in the temperature range 25e120 C. he second stages with an estimated mass loss of 32.62%, found in the temperature range of 120e320 C, which corresponds to a loss of one molecule of water, the guanine and fragment of iodide. The last decomposition step occur within the temperature range 320e600 C with an estimated
[Fe(Hdmg)2(Gua)Br]2H2O Fe(Hdmg)2(Gua) Fe(Hdmg)1.83
[Fe(Hdmg)2 (Gua)I]H2O [Fe(Hdmg)2 I0.83 Fe(Hdmg)1.5
mass loss at 40%, can be attributed to the liberation of iodide and partial elimination of Hdmg group (calculated mass loss 41.4%). The elimination of two molecules of dimethylglioxime occurs for all complexes after the elimination of HX (X ¼ Cl, Br or I) and the guanine [48]. The TG curves of the bromo complex are demonstrated in Fig. 14. 3.6. Antioxidant activity of complexes The antioxidant activity of the complexes was measured in terms of their hydrogen donating or radical scavenging ability by UVeVis spectrophotometer using the stable 2,2-diphenyl-1-picrylhydrazyl radical (DPPH). DPPH radicals are stable free radicals and in the presence of molecules capable of donating H atoms, its radical character is neutralized. This is visually noticeable as the color changes from purple to yellow. When complexes as antioxidant donate protons to this radical, the initial absorbance of DPPH solution decreases. The complex [Fe(Hdmg)2(Gua)I]3H2O and [Fe(Hdmg)2(Gua)Br]4H2O showed higher activity than the [Fe(Hdmg)2(Gua)Cl]6H2O one. This high activity suggests that the radical scavengers' activity of these complexes is probably due to an electron donor mechanism. The ranking of the antioxidant activity in the order presented in Table 6 is due to the effect of the elecctronegativity which increases in the order I < Br < Cl [49]. Several
Fig. 14. TG and DTG curves of [Fe(Hdmg)2(Gua)I]3H2O.
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422 Table 6 IC50 (mg/ml) of metal complexes. Materials
IC50(mg/ml)
A.A [Fe(Hdmg)2(Gua)I]3H2O (Complex 3) [Fe(Hdmg)2(Gua)Br] 4H2O (Complex 2) [Fe(Hdmg)2(Gua) Cl]6H2O (Complex 1)
0.13 0.34 4.5 12.8
research works have been carried out to study the antioxidant properties of iron complexes [50,51]. However, there is little known works on iron complexes with nitrogenous bases or with dimethylglyoxime. Minimum inhibitory concentrations IC50 of the tested compounds are illustrated in Fig. 15. The lower IC50 value (minimal inhibitory concentration) indicates a stronger ability of the tested compound to act as DPPH scavenger while the higher IC50 value indicates a lower scavenging activity of the scavangers. The complexes 3, 2 and 1 showed inhibitory concentration of 50% free radical IC50 of 0.34, 4.5 and 12.8 mg/ml respectively while ascorbic acid (AA positive control) presented a value of 0.13 mg/ml.
4. Conclusion In the present study, an attempt was made on to synthesize, characterize, and assesses the biological profile of the Iron (III) complexes of formula [Fe(Hdmg)2(Gua)X]nH2O have been synthesized and characterized by various physicochemical methods. For all the complexes an octahedral geometry were proposed, with two N-(dimethylglyoximato) ligands bonded to the metal in the equatorial plane. The two axial sites being occupied by a guanine and a halogen. The thermal behavior of the complexes, determined by thermogravimetry (TG) and differential thermally show that equatorial ligands are more stable that the ligands in axial position. In conclusion, the results obtained can help to confirm the structure of the compounds. We also found in the antioxidant activity that the lower the IC50 value, the greater the antioxidant activity of a
Fig. 15. The comparison of IC50 (mg/ml) of A.A and metal complexes.
421
compound. The complexes synthesized are endowed with an antioxidant capacity and capacity of scavenging the free radical DPPH. The results of this activity show that complex 3 has a similar antioxidant activity as A.A. Appendix A. Supplementary data Supplementary data to this article can be found online at https://doi.org/10.1016/j.molstruc.2019.02.108. References [1] T.M. Ross, S.M. Nevill, D.S. Innes, D.R. Turner, B. Moubaraki, K.S. Murray, The Royal Society of Chemistry Dalton,Trans. 7 (2010) 149e159. [2] S.V. Kryatov, E.V.R. Akimova, S. Schindle, Chem. Rev. 105 (2005) 2175e2226. [3] B. Nigovi c and N. Kujund zi c. 21 (2002) 1661e1666. [4] R. Singh, A. Banerjee, K. K Rajak, Inorg. Chim. Acta 363 (2005) 3131e3138. [5] M. Mohan, M. Kumar, Polyhedron 11 (1985) 1929e1939. [6] K. Burger, I. Ruff, F. Ruff, J. Inorg. Nucl. Chem. (1965) 179e190. [7] A. Chakravorty, Coord. Chem. Rev. 13 (1974) 1e46. [8] Y. Richter, B. Fischer, Nucleos Nucleot. Nucleic Acids 22 (2003) 57e1780. [9] A.J. Jeevagan, S.A. John, Anal. Biochem. 424 (2012) 21e26. [10] D.A. Ibrahim, A.M. Elmetwally, Eur. J. Med. Chem. 45 (2010) 1158e1166. [11] I. Machado, L.B. Marino, B. Demoro, G.A. Echeverria, O.E. Piro, C.Q. Leite, F.R. Pavan, D. Gambino, Eur. J. Med. Chem. 87 (2014) 267e273. [12] J.S. Lee, H.J. Kim, H. Park, Y.S. Lee, J. Nat. Prod. 65 (2002) 1367e1370. ~o, A.S. Reis, T.C. dos Santos, C.S. Coube, [13] L.L. Mensor, F.S. Menezes, G.G. Leita S.G. Leit~ ao, Phytother Res. 15 (2001) 127e130. [14] N. Raman, Y.R. pitchaikani, A. Kulandaisamy, Proc. Indiana Acad. Sci. 113 (2011) 183e189. [15] A. El-Dissouky, N.A. Al-Awadi, N.M. Shauib, A.B. Abbas, Spectrochim. Acta 67 (2007) 1072e1079. [16] M.S. Masoud, A.A. Soayed, A.E. Ali, Spectrochim. Acta, Part A 60 (2004) 1907e1915. [17] C.M. Mikulski, L. Mattucci, L. Wiss, N.M. Karayannis, Inorg. Chim. Acta 92 (1984) 275e278. [18] Y.Z. Voloshin, O.A. Varzatskii, A.V. Palchik, Polyhedron 17 (25) (1998) 4315e4326. [19] A. Adkhis, O. Benali-Baitich, M. Khan, G. Bouet, Org. Chem. 30 (2000) 1849. [20] S.C. Nayak, P.K. Doss, K. Sahoo, J. Anal. Appl. Pyrolysis 70 (2003) 699e709. [21] Y. Yang, J. Guan, P. Qiu, Q. Kan, Transition. Met. Chem. 35 (2010) 263e270. [22] N.T. Madhu, P.K. Radhakrishnan, M. Grunert, P. Weinberger, Linert W. Thermochimica, Acta 407 (2003) 73e84. [23] M. Sankaralingam, N. Saravanan, N. Anitha, E. Suresh, M. Palaniandavar, The Royal Society of Chemistry, Dalton Trans 43 (2014) 6828e6841. [24] M. Pascaly, M. Duda, F. Scweppe, K. Zurlinden, F.K. Muller, J, B. Krebs, Chem. Soc., Dalton Trans. (2001) 828e837. [25] A. El-Dissouk, N. Al-Awadi, N.M. Shauib, A.B. Abbas, Spectrochim. Acta, Part A 67 (2007) 1072e1079. [26] K.K. Narang, V.P. Singh, Synth. React. Inorg. Met. Org. Chem. 23 (1993) 971e989. [27] A.D. Becke, J. Chem. Phys. 98 (1993) 5648e5652. [28] A.D. Becke, J. Chem. Phys. 98 (1993) 1372e1377. [29] A.D. Becke, J. Chem. Phys. 97 (1992) 9173e9177. [30] C. Lee, W. Yang, R.G. Parr, Phys. Rev. B 37 (1988) 785e789. [31] T.H. Dunning Jr., P.J. Hay, H.F. Schaefer III (Eds.), Modern Theoretical Chemistry, vol. 3, Plenum, New York, 1976, p. 1. [32] P.J. Hay, W.R. Wadt, J. Chem. Phys. 82 (1985) 270e283. [33] P.J. Hay, W.R. Wadt, J. Chem. Phys. 82 (1985) 284e298. [34] P.J. Hay, W.R. Wadt, J. Chem. Phys. 82 (1985) 299e310. [35] F. Neese, The ORCA program system, WIREs Comput. Mol. Sci. 2 (2012) 73e78. [36] F. Neese, J. Chem. Phys. 127 (16) (2007) 164112e164119. [37] S.I. Gorelsky, A.B.P. Lever, J. Organomet. Chem. 635 (2001) 187e196. [38] D. Polyansky, D. Cabelli, J.T. Muckerman, E. Fujita, T.A. Koizumi, T. Fukushima, T. Wada, K. Tanaka, Angew Chem. Int. Ed. Engl. 46 (2007) 4169e4172. [39] C.E. Check, T.O. Faust, J.M. Bailey, B.J. Wright, T.M. Gilbert, L.S. Sunderlin, J. Phys. Chem. A 105 (2001) 8111e8116. [40] E.W.S. Caetano, U.L. Fulco, E.L. Albuquerque, A.H. de Lima Costa, S.N. Costa, A.M. Silva, F.A.M. Sales, V.N. Freire, J. Phys. Chem. Solids 121 (2018) 36e48. [41] D. Belkhir-Talbi, M. Makhloufi-Chebli, S. Terrachet- Bouaziz, D. HikemOukacha, N. Ghemmit, L. Ismaili, A. M.S Silva and M. Hamdi. 1179 (2019) 495-505. [42] T. Dhanalakshmi, E. Suresh, M. Palaniandavar, J. Chem. Soc. Dalton Trans. 39 (2009) 8317e8328. [43] Y.Z. Vdoshin, O.A. Varzatkii, A.V. Palchik, E.V. Polshin, J.A. Moletin, N.G. Strizhakova, Plyhedron 17 (26) (1998) 4315e4326. [44] M.T. Escot, P. Pouillen, P. Martinet, Electrochim. Acta (1981) 69e73. [45] M. Ocafrain, M. Devaud, M. Troupl, J. Perichon, Electrochim. Acta 42 (1997) 99e105. [46] N.R.S. Kumar, M. Nethaji, K.C. Patil, Polyhedron (1991) 365e371. [47] M.S. Masoud, A.M. Ramdan, G.M. El-Ashry, Thermochim. Acta 551 (2013)
422
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422
164e174. [48] E.S. Padial, D.C. Lazarte, E.B. Altaminaro, A. Castineira, R. Carrballo, rrez, Polyhedron 21 (2002) 1451e1457. J.N. Gutie [49] R.S. Priya, S. Balachandran, J. Daisy, P.V. Mohanan, Universal Journal of Physics and Application 3 (1) (2015) 6e16, https://doi.org/10.13189/
ujpa.2015.030102. [50] N.R. Perron, J.L. Brumghin, Cell Biochem. Biophys. 53 (2009) 75e100. [51] M. Leopoldini, N. Russo, S. Chiodo, M. Toscano, J. Agric. Food Chem. 45 (2000) 6343e6351.
Contents lists available at ScienceDirect
Journal of Molecular Structure journal homepage: http://www.elsevier.com/locate/molstruc
Synthesis, DFT/TD-DFT theoretical studies, experimental characterization, electrochemical and antioxidant activity of Fe(III) complexes of bis (dimethylglyoximato) guanine Lamia Abane-Merzouk a, Ahmed Adkhis a, *, Souhila Terrachet-Bouaziz b, Malika Makhloufi-Chebli a a b
Laboratoire LPCM, D epartement de Chimie, Facult e des Sciences, Universit e Mouloud Mammeri, 15000, TiziOuzou, Algeria Department of Chemistry, Faculty of Sciences, University Mohamed Bouguerra, Boumerdes, Algeria
a r t i c l e i n f o
a b s t r a c t
Article history: Received 13 January 2019 Received in revised form 21 February 2019 Accepted 26 February 2019 Available online 4 March 2019
A new series of Iron (III) complexes of general formula [Fe(Hdmg)2(Gua)X]nH2O where (Hdmg ¼ dimethylglyoximato monoanion, Gua ¼ Guanine, X ¼ Cl, Br or I) have been prepared and characterized by elemental analysis, conductivity, infrared spectra and electronic spectra. The molar conductance data confirm that all the complexes are no electrolytic. The mode of bonding and geometry of the Fe(III) complexes have been inferred spectral data, an octahedral geometry is ascribed for all the Fe(III) complexes. In these compounds, two dimethylglyoximato monoanions bond to the metal in the equatorial plan have been formed. The two axial sites being occupied by the guanine and the halogen moieties. The thermal properties of the Fe(III) complexes were investigated using TG and TDA. The thermal behavior of the compounds illustrates the general decomposition patterns of the complexes. The voltammetry experiments have been performed on the complexes to evaluate the redox process of Fe(III)/Fe(II) couple. These compounds were screened for their in-vitro antioxidant properties using 2,2diphenyl-1-picrylhydrazyl radical (DPPH .) free radical scavenging. The results obtained show that these complexes have a good antioxidant activity in comparison with ascorbic acid as positive control. © 2019 Published by Elsevier B.V.
Keywords: Dimethylglyoxime Iron(III) complexes Spectroscopic analysis Cyclic voltammetry Anti-oxidant activity Drug-likeness ADME analysis
1. Introduction Iron is the most abundant metal present at active sites of a large number of metalloenzymes [1]. The synthesis and characterization of iron complexes with different ligands play an important role in the coordination chemistry of iron because of their magnetic properties which exhibit spin states, low-spin and high-spin [2]. Furthermore it was found that some iron(III) complexes provide a useful structural and electronic model for the similarly coordinated iron(III) sites found in the hem iron enzymes [3,4]. On the other hand, oximes are well known for their high coordinating abilities and versatile behavior, but much remains to be learned about the type of structures that are formed as well as the factors that dictate structure. In general, dioximes coordinate with metal ions in four ways [5]. The most stable form corresponds to that shown in Fig. 1. This stability is due to the existence of the hydrogen bonds which
* Corresponding author. E-mail address: [email protected] (A. Adkhis). https://doi.org/10.1016/j.molstruc.2019.02.108 0022-2860/© 2019 Published by Elsevier B.V.
are formed between the hydroxyl groups. The dimethylglyoxime ligand (dmgH2) is a particulary interesting chelating oxime ligand [6]. The general characteristic of the structure of transition metal dioxime complexes are summarized early in the works of Chakravorty [7]. It is well known that the nucleotides acids vary in their ability to participate in many biological reactions forming metal complexes [8]. The diversity of research contribution has been affected by a large variety of studies concerning the chemical and biological aspects of the interactions between metal ion and nucleic acids and their constituent [9]. In recent years, interest in natural or synthetic antioxidants, in relation to therapeutic properties, has increased considerably. Antioxidants are the subject of numerous researches due to their importance in the health of pernicious diseases (cancer) and in the agri-alimo industry. Compounds containing pyrimidine and purine play a significant role in many biological systems, where both exist in nucleic acids, several vitamins, coenzymes and antibiotics [10]. These provide potential binding sites for metal ions, and any information on their coordinating properties is important as a means of understanding the role of the metal ions in biological systems.
414
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422
ethanolic solution (40 ml) and H2dmg (2.32 g, 20 mmol) dissolved in a hot ethanol (60 ml) was added powder guanine (1.51 g, 10 mmol). The reaction mixture is stirred and refluxed for approximately 4 h at a temperature of 50 C. The brown precipitated solid obtained was collected by filtration, washed with ethanol and diethylether. The solid was dried in air. Yield: 45%; Anal. Calc. for C13H33FeN9O11Cl: Fe, 7.43%; C, 27.02%; N, 21.82%; H, 7.27%. Found: Fe, 7.73%; C, 27.23%; N, 21.85%; H, 7.24%.
Fig. 1. Structure of dimethylglyoximato complexes M(Hdmg)2.
Metal ion-nucleic acid complexes and metal oxime complexes have extensive applications in the fields of medicine, photo and magnetic chemistry and electrochemistry. In addition, iron is needed by the bacteria and different iron complexes of bioactive ligands have shown anti-microbacterial activity [11]. Hence in this work we report the syntheses, characterization of iron(III) complexes, and their antioxidant activities.
2.2.2. Synthesis of [Fe(Hdmg)2(Gua) Br]4H2O (complex 2) To an aqueous solution (40 ml) of Fe(NO3)3.9H2O (4.04 g, 10 mmol) and NaBr (1.02 g, 10 mmol) was added dropwise a hot ethanolic solution (60 ml) of dimethylglyoxime (2.32 g, 20 mml). Then powdery guanine (1.51 g, 10 mmol) was added to the mixture. The resulting mixture is left stirring under reflux at 50 C for 3 h. The brown precipitate obtained was isolated by filtration and washed with ethanol and diethylether. The solid was dried at room temperature. Yield: 37%; Anal. Calc. for C13H29FeN9O9Br: Fe, 8.39%; C, 26.63%; N, 21.5%; H, 4.09%. Found: Fe, 8.47%; C, 26.57%; N, 21.20%; H, 4.02%.
2. Experimental
2.2.3. Synthesis of [Fe(Hdmg)2(Gua)I]3H2O (complex 3) This compound was synthesized using the same procedure as described above for complex II, starting from KI (1.66 g, 10 mmol) instead of NaBr. Yield: 42%; Anal. Calc. for C13H27FeN9O8I: Fe, 9.36%; C, 25.38%; N, 20.49%; H, 3.57%. Found: Fe, 9.78%; C, 25.94%; N, 20.04%; H, 3.37%.
2.1. Materials and methods
2.3. Antioxidant activity
The chemicals reagents and solvents (Fluka products) used were of analytical grade and were used without further purification. Melting points were determined on a Stuart scientific SPM3 apparatus fitted with a microscope and are uncorrected. The infrared spectra were recorded in the region 4000e400 cm1 on a BRUKER TENSOR 27 IR spectrophotometer. Electronic spectra were measured on a thermo Scientific EVOLUTION 220 ultravioletevisible spectrophotometer (in DMSO solution); measurements were made from 200 to 800 nm. The elemental microanalysis (C, H, N) was carried out at the Service of analysis and characterization of CNRS (Poitier France. The molar Conductance values were obtained for 10 3M in DMSO solution at 25 ± 1 C determined using melting point meter MPM-H2. All electrochemical experiments were carried out using a potentiostatgalvanostat EGG-273 A controlled with Power-suite software. The electrochemical cell was equipped with a modified carbon past disk as the working electrode, a platinum electrode as counter electrode and silver/silver chloride (Ag/AgCl) electrode as the reference electrode. The electrolytic bath consisted of the complexes in DMSO solution (103 M) and sodium perchlorate (NaClO4 101 M) was used as a supporting electrolyte to increase the conductivity of the bath. The electrochemical characterization was performed using the cyclic voltammetry (CV) with a scan rate of 50 mV/s. TG/DTA/ EDS techniques were obtained with NETZSCH STA 409 PC/PG. The heating rate of 10 C/mn was used between 25 C and 600 C. This theory uses Becke's hybrid three-parameter method and the Lee-Yang-Parr correlation function (B3LYP) associated with the use of ORCA V4 software, which has proved very effective for a large number of organic and inorganic systems.
The antioxidant activity (free radical scavenging activity) of the synthesized H2dmg metal complexes was evaluated using the 2,2diphenyl-1-picrylhydrazyl free radical (DPPH.) scavenging assay [12,13]. DPPH. is a stable free radical which has been widely accepted as a tool for estimating the free radical scavenging activities of potential antioxidants (Fig. 2) [14]. The free radical scavenging ability of H2dmg metal complexes was studied spectrophotometrically using DPPH radical scavenging technique. Radical scavenging activities of all complexes were determined by the interacting ability of compounds with stable free radical DPPH. The degree of discoloration indicated the scavenging potential of the antioxidant compounds in terms hydrogen donating ability. DPPH. solution was prepared by dissolving DPPH. in ethanol to give a concentration of 4 mg/100 mL1. The metal complexes were dissolved in DMSO to obtain a solution of 101 M. The concentration of test compounds was diluted with DMSO to get final concentrations 0.05, 0.025 et 0.0125 mol/l for all the compounds, but the standard were further diluted to give concentration solutions of 0.1, 0.05, 0.025 et 0.0125, 0.00625, 0.003125, 0.0015625 mol/l. 40 mL of each tested concentration of each complex was added to each well separately in diplicate, and then 2 mL of DPPH solution was added. Negative control wells were loaded with 40 mL of DMSO and 2 mL of DPPH. solution each. After vortexing, the mixtures were incubated at room temperature for 1 h in darkness at 25 C, and
2.2. Synthesis of iron(III) complexes 2.2.1. Synthesis of [Fe(Hdmg)2(Gua)Cl]6H2O (complex 1) To a mixture of FeCl3.6H2O (2.70 g, 10 mmol) dissolved in aqua
Fig. 2. Structure of the DPPH. radical and its DPPH2 reduction product.
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422
then the absorbance of these compounds at different concentrations was recorded at 517 nm. The reduction of the DPPH radical was measured by monitoring continuously the decrease of absorption at 517 nm. DPPH scavenging effect was calculated as percentage of DPPH discoloration using the equation:
RSA ð%Þ ¼ ½ðAc AsÞ=Ac x 100 where Ac is the absorbance of the control (absorbance of DPPH. ethanol solution without sample), and As is the absorbance of the tested compound after 60 min incubation. Minimum inhibitory concentrations IC50 of the tested compounds are determined from RSA (%) against concentration graph. 3. Results and discussion The complexes were obtained as powders and attempts to obtain a single crystal suitable for X-ray determination were unsuccessful. Their physical properties and analytical data are summarized in Table 1 and their empirical formulas were obtained based on elemental analysis and TGA measurements. The complexes are insoluble in water and soluble in DMF and DMSO. The results of elemental analysis suggested that the complexes can be formulated as [Fe(Hdmg)2(Gua)X]nH2O, showing that dimethylglyoximato group as a monoanion. The melting point of all complexes exceeds 360 C indicating that all complexes are stable in air and can be preserved at room temperature. The molar conductance values of all complexes in DMSO at 103 M supporting the non-electrolyte nature of the all iron complexes. 3.1. UVevisible spectroscopic study on the effect of the stoichiometric amount of the metal on the ligand To optimize the complexation, a study was carried out to determine the molar ratio M/Ls required for the synthesis of the complexes. The solution of the ligands and that of the metal were prepared in Ethanol-water (V/V) at concentrations of 105 M for the ligands and 0.125 105, 0.25 105, 0.5 105, 0.75 105 and 1 105 M for the Fe(NO3)3.9H2O. The absorption spectrums of free ligands H2dmg and guanine and of the mixture H2dmg-guanine (2:1) have been studied (Fig. 3AeC). In the case of the mixture H2dmg H2dmg-guanine, two bands were observed and the intense long-wavelength band situated around 231 nm was attributed to the p - p* transition. In the O present work, addition of Fe(NO)3.9H 3 2 to a solution of H2dmg-guanine 5 (2:1) (10 M) in ethanol-H2O (V/V) induced drastic changes in the absorption spectrum. A strong hypochromic effect on the wavelength band at 231 nm (Fig. 4, curves a to f) was observed. The band around 302 nm underwent a hypechromic effect. The absorbance was analysed as a function of cation concentration (Fig. 5). Since the shape of the absorption spectrum varied strongly in the presence of cations, the absorbance was recorded at two different wavelengths chosen in order to obtain maximum
415
information. The calculation was performed simultaneously on the two bands obtained by absorption spectroscopy. Good fits were obtained, as displayed in Fig. 5. The results of Fig. 4 show that from an equimolar ratio metal/ guanine and 2 equivalents of H2dmg, the absorbance becomes important. The time factor on metal (1eq)-ligands (H2dmg-guanine (2:1 eq)) coordination (1: 2:1) is also studied. Hypechromic effect on the wavelength bands at 231 nm and 302 nm is observed (Fig. 6). 3.2. Characterization of the ligands and the complexes 3.2.1. Infrared spectra In the absence of X-ray structure determination, infrared spectra are the most suitable technique to determine the coordinating atoms of the ligands by comparing the free ligand spectra with those of the complex. The main vibrations of the ligands and the complexes are presented in Table 2. The IR spectra for complexes 1, 2 and 3 are rather similar; the discussion is confined to the most important vibrations. The IR spectra of the complexes were compared with the spectra of H2dmg and guanine to establish the coordination modes of the ligands in the complexes. The spectra of all the complexes show bands in the region 3335-3338 cm1 due to the NH2 stretching vibration of guanine which is shifted to higher frequency that encountered in the free guanine (3322 cm1), the spectrum of the chloro complex is given by way of example in Fig. 7. This considerable shifts in ʋ(NH2) clearly indicate the involvement of NH2 of guanine in metal coordination [15]. The guanine exhibit a ʋ(CH) stretching vibration bands at 2991 cm1 and 2900 cm1. In the spectra of the complexes, these bands are located at around 2995 cm1 and 2902 cm1. The bands are apparently insensitive to the complexation [16]. Characteristic IR strong bands of free guanine corresponding to ʋ(C¼O) observed at 1693 and 1670 cm1, are shown in spectra of the complexes at 1690e1699 cm1 and 16611667 cm1 region. The little negative shifts indicate that the oxygen atom of carbonyl function is not participating to the metal coordination [17]. An additional band which appears at 3445, 3450 and 3465 cm1 respectively in the spectra of all the complexes is correlated with the ʋ(OH) vibration of water hydratation molecules. In the infrared spectra of H2dmg, the sharp band of medium intensity occurring at 1447 cm1 is attributed to y (C¼N) of the oxime. In the complexes, this band is shifted to higher frequency (1557 cm1) [18]. This suggests that dimethylglyoxime is coordinated to the metal ion through the nitrogen atom of the oxime. In all the complexes, the strong as. Well as sharp band between 1208 and 1206 cm1 is due to the NeO stretching vibration. The metal sensitive as sharp band at 772779 cm1 region may be attributed to the C¼NeO deformation vibration [19]. The peak around 1363 cm1 in the spectra of H2dmg is assigned to the deformation of -CH3 groups; this peak is observed in 1378e1387 cm1 region in the spectra of the compounds [20]. The characteristic stretching vibrations belonging to y(OH) of
Table 1 Analytical data of the complexes. M.C (U1.cm2. mol1)
Compound Formula
Color
Mol.Wt (g/mol)
Yield (%)
M.P (C )
Anal-Found/Cal (%) C
H
N
Fe
Fe[(Hdmg)2(Gua)Cl]6H2O
Brown
577.66
45
>360
Fe[(Hdmg)2Gua)Br]4H2O
Brown- Russ
586.11
37
>360
[F Fe[(Hdmg)2 (Gua)I]3H2O
Brown
615.11
42
>360
27.23 27.02 26.57 26.63 25.94 25.38
7.24 7.27 4.02 4.09 3.37 3.57
21.85 21.82 21.2 21.5 20.04 20.49
7.73 7.43 8.47 8.39 9.78 9.36
4.14 4.28 5.27
416
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422
Fig. 4. Absorption spectrum of the ligands (105 M) in ethanol in the absence (curve a) O Fe(NO) .9HO and presence of Fe(NO)3.9H 3 2 . Curves b to f: effect 33 2 addition. From bottom to top: Fe(NO) 9HO 5 5 5 5 and 2.105 M. 33 2 : 0.125.10 ; 0.25.10 ; 0.5.10 ; 10
Fig. 5. Fitted spectrophotometric data for the ligands vs. cation concentration. Variation of absorbance: l ¼ 231 and 304 nm.
the complexes 1, 2 and 3 is the absorption at 509e506 and 503 cm1 respectively; this was assigned to the (FeeN) vibration. These bands are not found in the spectrum of the ligand [22]. The variation in the position of these bands follows the order in spectrochemical series: I < Br < Cl. The representative IR spectra of [Fe(Hdmg)2(Gua)Cl]4H2O complex is shown in Figure and the other spectra are given in the Supplementary material section.
Fig. 3. (A) Absorption spectrum of H2dmg (105 M); (B) Absorption spectrum of guanine (105 M); (C) Absorption spectrum of H2dmg eguanine(2:1) (105 M).
oxime is observed at 3205 cm1 in dimethylglyoxime, this band is expected to disappear upon complex formation. The band in the region 947-951 cm1 may be attributed to deformation vibration of OH in H2dmg moiety [21]. The only new feature in the IR spectra of
3.2.2. Electronic spectra The electronic absorption spectrum is often very useful on the evaluation of structural investigation. The electronic absorption spectra of Fe(III) complexes was recorded at room temperature using DMSO as solvent. The spectra of H2dmg, guanine and its complexes with tentative assignments are presented in Table 3. The H2dmg exhibit two absorption bands in the UV, the first band with a strong intensity appear at 227 nm, the second lower intensity band appear at 313 nm (Fig. 3A). These bands can be attributed to p / p* transition of the mimic (C¼N) function present
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422
417
Fig. 7. Experimental FT-IR spectra of [Fe(Hdmg)2(Gua)Cl]4H2O. Fig. 6. Absorption spectrum of the ligands (105 M) in ethanol in the presence of Fe(NO) 9HO 5 M) along the time. 33 2 (10
in the ligand moiety. The spectra of guanine display in the UV an intense band at 273 nm which could be ascribbed to p / p* transition (Fig. 3B). The electronic spectra of the three complexes show intense bands at 289, 304 and 296 nm respectively, which could be assigned to p / p* transition. The position of these bands is shifted to higher wavelength (lower energy) compared to that observed in the free ligand (for example UVevisible spectra of complex 1, Fig. 8). The chloro, bromo and iodo dimethylglyoximato Fe(III) complexes exhibit similar spectra in the UV and in the visible region. While the spectral features observed in the range 322e376 nm may be assigned to X Fe(III) (X ¼ Cl, Br or I) charge transfer transition [23]. These values show that the general trend observed is shifting of the L/M transition bands in the complexes increases with the growth of electronegativity of the axial ligand: I < Br < Cl [24]. In addition, the very relatively low intense band appear at around 473 nm assignable to the spin allowed electronic transition corresponding to 6A1g / 4T1g (G) transition consistent with octahedral Fe(III) complexes [25,26]. The results obtained describe a pseudo-octahedral structure for all iron complexes (Fig. 9), where the two dimethylglyoxime monoanions are in the equatorial plane. The guanine and X (halogen) occupy the two axial coordination sites.
Table 3 Electronic spectra data of iron (III) Complexes. Complexes
l(nm)
ῡ(cm1)
Electronic transition
H2dmg Guanine
227 313 273 289 376 486 304 328 477 296 322 356 473
44053 31949 33630 34602 26596 20576 32894 30487 20964 33784 31056 28090 21141
p - p* p - p* p - p* p - p*
Complex 1
Complex 2
Complex 3
C.T. Cl/Fe(III) Transition d-d p - p* C.T. Br/Fe(III) Transition d-d p - p* C.T. I /Fe(III) Transition d-d
3.3. Complex structures and theoretical studies In order to determine theoretically the geometry of the complexes and make a comparison with the experimental observations, quantum chemical optimisations of the I. R and UVeVISIBLE have been carried out. The geometries of the synthesized complexes were optimized with density functional theory (DFT) using Becke's three parameters hybrid method and the Lee-Yang-Parr correlation functional (B3LYP) [27e30] combined with LANL2DZ [31e34] using ORCA v4
Table 2 Relevant infrared spectra bands for the ligands and Complexes (cm1). Complexes
y(OH)
H2dmg Guanine
3205
Complex 1
e
Complex 2
e
Complex 3
e
y
as(NH2) s
y (NH2)
3320 3116 3335 3115 3335 3109 3338 3112
n (CH)
y(C¼O)
n (CN)
n(NO)
d (NO)
n (FeeN)
e 2991 2905 2993 2904 e 2903 2995 2902
e 1693 1670 1694 1669 1699 1661 1690 1667
1447 e e 1557
1144 e e 1208
750 e e 776
e e e 509
1557
1206
772
506
1557
1208
779
503
418
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422
Fig. 10. Theoretical electronic Spectra of complex 1 (X ¼ Cl).
Fig. 8. UVevisible spectra of the complex [Fe(Hdmg)2(Gua)Cl]6H2O in DMSO.
Fig. 9. Suggested structure of the complexes. (X ¼ Cl, Br or I). Fig. 11. Theoretical FT-IR spectra of complex 1 (X ¼ Cl).
software [35,36]. Infrared spectra were obtained at the same level of theory combined with Time-Dependent Density Functional Theory (TD-DFT) for UVevis absorption. The calculated UVevis absorption is in agreement with the experimental results. It is worth noting that previous studies were demonstrated the capability of chosen method and basis set to predict the electronic properties of metallic complexes quite reasonably [37e41]. For example, the calculated absorption spectrum of the complex [Fe(Hdmg)2(Gua)Cl]4H2O (Fig. 10) shows in the ultraviolet domain a band around 320 nm due to the transition p / p *. Another band located at 430 nm corresponds to the metalligand charge transfer. A band appears in the visible which correspond to the transition 6A1g / 4T1g. The ligands coordination sites which are involved in bonding with the metal ions have been determined by careful comparison of the infrared absorption spectra of the complexes with that of the parent ligand. The vibrational assignments were carried out with support DFT calculations. The theoretical FT-IR spectra of [Fe(Hdmg)2(Gua)Cl]4H2O is given in Fig. 11 and compared with the experimental spectra. The IR spectra indicate that the ligands are coordinated with the metal, resulting in neutral species. Coordination occurs through the N-atoms from the H2dmg and Guanine moiety and from the
halogen. On basis of the elemental analysis, electronic spectra and theoretical studies, the complexes were formulated as Fe [(Hdmg)2(Gua)Cl]6H2O, Fe[(Hdmg)2Gua)Br]4H2O and Fe[(Hdmg)2 (Gua)I]3H2O (Fig. 12). 3.4. Electrochemical: cyclic voltammetry study The application of electrochemical methods to the study of the coordination of metal ion provides a useful complement to the previously used methods of investigation, such as UVevisible spectroscopy. The electrochemical behaviours of the ligands and complexes were examined by cyclic voltammetry. DMSO solution (103 M) containing 101 M sodium perchlorate monohydrate (NaClO4) as supporting electrolyte was used. The potential were scanned in the range [-1,4 V, þ1,4 V] at a scan rate of 50 mV s1. The cyclic voltammogram of guanine display two cathodic waves at 0.78 V and 0.47 V of the molecule, one anodic waves is observed at þ0.75 V around, due probably to the oxidation of guanine. In the voltammogram of H2dmg, one anodic wave is located at þ0.95 V
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422
419
Table 4 Cyclic voltametry data for the iron(III) complexes. Complexes Complexes
Couple red/ox
Epc[V]
Epa[V]
DEP[mV]
ipa/ipc
Complex 1
Fe(III/II) Fe(III/II)
Complex 3
Fe(III/II)
0.52 þ0.26 0.58 þ0.23 0.57 þ0.57
500 370 120 530 150 e
0.38
Complex 2
1.02 0.11 0.70 0.30 0.72 e
0.27 0.42 0.10 e
assigned to the oxidation of the oxime group [42]. The reduction potential and the corresponding values summarized in Table 4, correspond to one electron irreversible metal centered Fe(III)/Fe(II) redox processes. The representative cyclic voltammogram of [Fe(Hdmg)2(Gua)Br] 4H20 complex is shown in Fig. 13. The complexes exhibited irreversible one electron response (DE ˃ 100 m V) corresponding to the couple Fe(III)/(II) [43]. This observation can be explained by the break of the bonding Fe (III)-X lasting the reduction of the complexes [Fe(III)(Hdmg)2GuaX]. So the axial ligand (X) is lost, and the formation of pentacoordined complex of Fe(II) according to the reaction (*) can be envisaged without the axial ligand X. [Fe / [Fe(II)(Hdmg)2 (Gua)] þ X (X ¼ Cl, Br, (III)(Hdmg)2(Gua)X] þ e I).(*) [44]. Values in Table 4 show that for all the complexes, the current ratio of the direct and reverse reduction waves is less then unity. This indicates the low chemical stabilities of this species in time scale. The heigh stability of the reduced forms of the complexes is also confirmed by the high DEp value which is compelling evidence for the irreversibility of electron transfer reaction. The cyclic voltammetry of complexes I and II presents also two defined cathodic waves at 0.11 V and 0.30 V, respectively and two anodic waves at þ0.26 V and þ0.23 V. The redox processes must be considered irreversible. In the other hand, defined oxidation processes in the cyclic voltammetry of the complexes II and III is observed at positive potentials þ0.56 and þ 0.57 V respectively, these potentials can be attributed to oxime function in the complexes. 3.5. Thermal study The experimental data corresponding to the chloro, bromo and iodo complexes are listed in Table 5. The thermograms indicate that all complexes involved water molecules [45]. The three complexes show decomposition pattern of three stages.
Fig. 12. Optimized Geometry for complexes 1 (A); 2 (B) and 3 (C). Fig. 13. Cyclic voltammetry curve of [Fe(Hdmg)2(Gua)Br]4H2O.
420
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422
Table 5 Thermal decomposition data of the Fe (III) Complexes. Compound
Temperature Range (C)
Weight loss (Obs)
Weight loss (Cal)
Species formed
Complex 1
25e120 120e300 300e600 25e160 160e380 380e600 25e120 120e320 320e600
3.13 20.94 62.16 6.44 22.79 38.32 5.74 32.62 40.00
3.11 19.24 62.59 6.14 21.06 38.32 5.85 32.68 41.60
[Fe(Hdmg)2(Gua)Cl]5H2O [Fe(Hdmg)2(Gua)Cl0.56] Fe(Hdmg)1.75
Complex 2
Complex 3
For the chloro complex, the first steps with estimated mass loss of 3.13%, found within the temperature range 25e120 C, corresponding to loss of uncoordinated water molecule. He second steps with estimated weight loss of 20.94% within the temperature rang 120e300 C, can be attributed to the liberation of five latter water molecules and approximately moiety of the chloride in the structure [46,47]. The third step is started at 300 C and ends at 600 C, the corresponding mass loss (62.12%) is due the decomposition of the guanine molecule and a fragment of Hdmg. The mass loss in the third stage is in agreement with the masse calculated loss of 62.59%. The bromo complex exhibit one weight loss step over the temperature ranges 25e120 C, due to the elimination of the two water molecules of hydration. The second step which corresponds to an endothermic peak corresponds to the elimination of two water molecules hydration and the bromide. The third thermal reaction is endothermic within 380e600 C range. The mass loss in this stage corresponds approximately to the guanine and 0.21 Hdmg. In iodo complex, the single step weight loss corresponds to simultaneous disappearance of two water molecules hydration in the temperature range 25e120 C. he second stages with an estimated mass loss of 32.62%, found in the temperature range of 120e320 C, which corresponds to a loss of one molecule of water, the guanine and fragment of iodide. The last decomposition step occur within the temperature range 320e600 C with an estimated
[Fe(Hdmg)2(Gua)Br]2H2O Fe(Hdmg)2(Gua) Fe(Hdmg)1.83
[Fe(Hdmg)2 (Gua)I]H2O [Fe(Hdmg)2 I0.83 Fe(Hdmg)1.5
mass loss at 40%, can be attributed to the liberation of iodide and partial elimination of Hdmg group (calculated mass loss 41.4%). The elimination of two molecules of dimethylglioxime occurs for all complexes after the elimination of HX (X ¼ Cl, Br or I) and the guanine [48]. The TG curves of the bromo complex are demonstrated in Fig. 14. 3.6. Antioxidant activity of complexes The antioxidant activity of the complexes was measured in terms of their hydrogen donating or radical scavenging ability by UVeVis spectrophotometer using the stable 2,2-diphenyl-1-picrylhydrazyl radical (DPPH). DPPH radicals are stable free radicals and in the presence of molecules capable of donating H atoms, its radical character is neutralized. This is visually noticeable as the color changes from purple to yellow. When complexes as antioxidant donate protons to this radical, the initial absorbance of DPPH solution decreases. The complex [Fe(Hdmg)2(Gua)I]3H2O and [Fe(Hdmg)2(Gua)Br]4H2O showed higher activity than the [Fe(Hdmg)2(Gua)Cl]6H2O one. This high activity suggests that the radical scavengers' activity of these complexes is probably due to an electron donor mechanism. The ranking of the antioxidant activity in the order presented in Table 6 is due to the effect of the elecctronegativity which increases in the order I < Br < Cl [49]. Several
Fig. 14. TG and DTG curves of [Fe(Hdmg)2(Gua)I]3H2O.
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422 Table 6 IC50 (mg/ml) of metal complexes. Materials
IC50(mg/ml)
A.A [Fe(Hdmg)2(Gua)I]3H2O (Complex 3) [Fe(Hdmg)2(Gua)Br] 4H2O (Complex 2) [Fe(Hdmg)2(Gua) Cl]6H2O (Complex 1)
0.13 0.34 4.5 12.8
research works have been carried out to study the antioxidant properties of iron complexes [50,51]. However, there is little known works on iron complexes with nitrogenous bases or with dimethylglyoxime. Minimum inhibitory concentrations IC50 of the tested compounds are illustrated in Fig. 15. The lower IC50 value (minimal inhibitory concentration) indicates a stronger ability of the tested compound to act as DPPH scavenger while the higher IC50 value indicates a lower scavenging activity of the scavangers. The complexes 3, 2 and 1 showed inhibitory concentration of 50% free radical IC50 of 0.34, 4.5 and 12.8 mg/ml respectively while ascorbic acid (AA positive control) presented a value of 0.13 mg/ml.
4. Conclusion In the present study, an attempt was made on to synthesize, characterize, and assesses the biological profile of the Iron (III) complexes of formula [Fe(Hdmg)2(Gua)X]nH2O have been synthesized and characterized by various physicochemical methods. For all the complexes an octahedral geometry were proposed, with two N-(dimethylglyoximato) ligands bonded to the metal in the equatorial plane. The two axial sites being occupied by a guanine and a halogen. The thermal behavior of the complexes, determined by thermogravimetry (TG) and differential thermally show that equatorial ligands are more stable that the ligands in axial position. In conclusion, the results obtained can help to confirm the structure of the compounds. We also found in the antioxidant activity that the lower the IC50 value, the greater the antioxidant activity of a
Fig. 15. The comparison of IC50 (mg/ml) of A.A and metal complexes.
421
compound. The complexes synthesized are endowed with an antioxidant capacity and capacity of scavenging the free radical DPPH. The results of this activity show that complex 3 has a similar antioxidant activity as A.A. Appendix A. Supplementary data Supplementary data to this article can be found online at https://doi.org/10.1016/j.molstruc.2019.02.108. References [1] T.M. Ross, S.M. Nevill, D.S. Innes, D.R. Turner, B. Moubaraki, K.S. Murray, The Royal Society of Chemistry Dalton,Trans. 7 (2010) 149e159. [2] S.V. Kryatov, E.V.R. Akimova, S. Schindle, Chem. Rev. 105 (2005) 2175e2226. [3] B. Nigovi c and N. Kujund zi c. 21 (2002) 1661e1666. [4] R. Singh, A. Banerjee, K. K Rajak, Inorg. Chim. Acta 363 (2005) 3131e3138. [5] M. Mohan, M. Kumar, Polyhedron 11 (1985) 1929e1939. [6] K. Burger, I. Ruff, F. Ruff, J. Inorg. Nucl. Chem. (1965) 179e190. [7] A. Chakravorty, Coord. Chem. Rev. 13 (1974) 1e46. [8] Y. Richter, B. Fischer, Nucleos Nucleot. Nucleic Acids 22 (2003) 57e1780. [9] A.J. Jeevagan, S.A. John, Anal. Biochem. 424 (2012) 21e26. [10] D.A. Ibrahim, A.M. Elmetwally, Eur. J. Med. Chem. 45 (2010) 1158e1166. [11] I. Machado, L.B. Marino, B. Demoro, G.A. Echeverria, O.E. Piro, C.Q. Leite, F.R. Pavan, D. Gambino, Eur. J. Med. Chem. 87 (2014) 267e273. [12] J.S. Lee, H.J. Kim, H. Park, Y.S. Lee, J. Nat. Prod. 65 (2002) 1367e1370. ~o, A.S. Reis, T.C. dos Santos, C.S. Coube, [13] L.L. Mensor, F.S. Menezes, G.G. Leita S.G. Leit~ ao, Phytother Res. 15 (2001) 127e130. [14] N. Raman, Y.R. pitchaikani, A. Kulandaisamy, Proc. Indiana Acad. Sci. 113 (2011) 183e189. [15] A. El-Dissouky, N.A. Al-Awadi, N.M. Shauib, A.B. Abbas, Spectrochim. Acta 67 (2007) 1072e1079. [16] M.S. Masoud, A.A. Soayed, A.E. Ali, Spectrochim. Acta, Part A 60 (2004) 1907e1915. [17] C.M. Mikulski, L. Mattucci, L. Wiss, N.M. Karayannis, Inorg. Chim. Acta 92 (1984) 275e278. [18] Y.Z. Voloshin, O.A. Varzatskii, A.V. Palchik, Polyhedron 17 (25) (1998) 4315e4326. [19] A. Adkhis, O. Benali-Baitich, M. Khan, G. Bouet, Org. Chem. 30 (2000) 1849. [20] S.C. Nayak, P.K. Doss, K. Sahoo, J. Anal. Appl. Pyrolysis 70 (2003) 699e709. [21] Y. Yang, J. Guan, P. Qiu, Q. Kan, Transition. Met. Chem. 35 (2010) 263e270. [22] N.T. Madhu, P.K. Radhakrishnan, M. Grunert, P. Weinberger, Linert W. Thermochimica, Acta 407 (2003) 73e84. [23] M. Sankaralingam, N. Saravanan, N. Anitha, E. Suresh, M. Palaniandavar, The Royal Society of Chemistry, Dalton Trans 43 (2014) 6828e6841. [24] M. Pascaly, M. Duda, F. Scweppe, K. Zurlinden, F.K. Muller, J, B. Krebs, Chem. Soc., Dalton Trans. (2001) 828e837. [25] A. El-Dissouk, N. Al-Awadi, N.M. Shauib, A.B. Abbas, Spectrochim. Acta, Part A 67 (2007) 1072e1079. [26] K.K. Narang, V.P. Singh, Synth. React. Inorg. Met. Org. Chem. 23 (1993) 971e989. [27] A.D. Becke, J. Chem. Phys. 98 (1993) 5648e5652. [28] A.D. Becke, J. Chem. Phys. 98 (1993) 1372e1377. [29] A.D. Becke, J. Chem. Phys. 97 (1992) 9173e9177. [30] C. Lee, W. Yang, R.G. Parr, Phys. Rev. B 37 (1988) 785e789. [31] T.H. Dunning Jr., P.J. Hay, H.F. Schaefer III (Eds.), Modern Theoretical Chemistry, vol. 3, Plenum, New York, 1976, p. 1. [32] P.J. Hay, W.R. Wadt, J. Chem. Phys. 82 (1985) 270e283. [33] P.J. Hay, W.R. Wadt, J. Chem. Phys. 82 (1985) 284e298. [34] P.J. Hay, W.R. Wadt, J. Chem. Phys. 82 (1985) 299e310. [35] F. Neese, The ORCA program system, WIREs Comput. Mol. Sci. 2 (2012) 73e78. [36] F. Neese, J. Chem. Phys. 127 (16) (2007) 164112e164119. [37] S.I. Gorelsky, A.B.P. Lever, J. Organomet. Chem. 635 (2001) 187e196. [38] D. Polyansky, D. Cabelli, J.T. Muckerman, E. Fujita, T.A. Koizumi, T. Fukushima, T. Wada, K. Tanaka, Angew Chem. Int. Ed. Engl. 46 (2007) 4169e4172. [39] C.E. Check, T.O. Faust, J.M. Bailey, B.J. Wright, T.M. Gilbert, L.S. Sunderlin, J. Phys. Chem. A 105 (2001) 8111e8116. [40] E.W.S. Caetano, U.L. Fulco, E.L. Albuquerque, A.H. de Lima Costa, S.N. Costa, A.M. Silva, F.A.M. Sales, V.N. Freire, J. Phys. Chem. Solids 121 (2018) 36e48. [41] D. Belkhir-Talbi, M. Makhloufi-Chebli, S. Terrachet- Bouaziz, D. HikemOukacha, N. Ghemmit, L. Ismaili, A. M.S Silva and M. Hamdi. 1179 (2019) 495-505. [42] T. Dhanalakshmi, E. Suresh, M. Palaniandavar, J. Chem. Soc. Dalton Trans. 39 (2009) 8317e8328. [43] Y.Z. Vdoshin, O.A. Varzatkii, A.V. Palchik, E.V. Polshin, J.A. Moletin, N.G. Strizhakova, Plyhedron 17 (26) (1998) 4315e4326. [44] M.T. Escot, P. Pouillen, P. Martinet, Electrochim. Acta (1981) 69e73. [45] M. Ocafrain, M. Devaud, M. Troupl, J. Perichon, Electrochim. Acta 42 (1997) 99e105. [46] N.R.S. Kumar, M. Nethaji, K.C. Patil, Polyhedron (1991) 365e371. [47] M.S. Masoud, A.M. Ramdan, G.M. El-Ashry, Thermochim. Acta 551 (2013)
422
L. Abane-Merzouk et al. / Journal of Molecular Structure 1186 (2019) 413e422
164e174. [48] E.S. Padial, D.C. Lazarte, E.B. Altaminaro, A. Castineira, R. Carrballo, rrez, Polyhedron 21 (2002) 1451e1457. J.N. Gutie [49] R.S. Priya, S. Balachandran, J. Daisy, P.V. Mohanan, Universal Journal of Physics and Application 3 (1) (2015) 6e16, https://doi.org/10.13189/
ujpa.2015.030102. [50] N.R. Perron, J.L. Brumghin, Cell Biochem. Biophys. 53 (2009) 75e100. [51] M. Leopoldini, N. Russo, S. Chiodo, M. Toscano, J. Agric. Food Chem. 45 (2000) 6343e6351.