Iron Complexes Containing Electrochemically Active Diazocycle-bis ( di-tert-butyl-phenol ) Ligands

Quatro ligantes N,O-doadores contendo unidades centrais diazocíclicas e grupos bis(diterc-butilfenol) foram empregados na síntese de complexos de ferro(III), resultando em quatro complexos mononucleares e um binuclear. Os ligantes apresentados neste estudo diferem entre si na unidade diazocíclica, sendo elas: piperazina (H2L1), diazepam/homopiperazina (H2L2), hexaidropirimidina (H2L3) ou hexaidropirimidin-5-ol (H3L4). As estruturas moleculares dos complexos [FeL2Cl], 2, e [Fe2(L4)(HL4)Cl], 4, foram elucidadas por difratometria de raios X de monocristal. Estudos eletroquímicos mostram que, além de processos redox centrados no metal, os complexos apresentam processos redox atribuídos aos ligantes. Estudos coulométricos acoplados a espectroscopia eletrônica no UV-Vis confirmam a formação da espécie radicalar para o complexo 2, enquanto dados de espectroscopia de ressonância paramagnética eletrônica (EPR) mostram a formação do radical para os complexos 2, 4 e 5. De relevância é a observação de que o ligante que sofre oxidação em menor potencial eletroquímico é aquele cujo centro metálico apresenta a menor acidez de Lewis.


Introduction
The synthesis of coordination compounds employing redox active ligands has been considered of interest in several fields of inorganic chemistry, e.g., coordination chemistry, 1 catalysis, 2 and bioinorganic chemistry. 3Examples of such ligands are α-diimines, 4,5 tert-butylphenol, [6][7][8] and porphyrins. 9The presence of electrochemically active ligands has also been documented in natural systems.0][11] It is also very well known that phenoxyl radical is formed in the galactose oxidase catalytic cicle, 12,13 as well as it has been observed in ribonucleotide reductase. 9s presented by Kaim, 14 the research involving redox active ligands poses several interests, including (i) the identification and the establishment of the non-innocent behavior; (ii) design and development of new ligands and their metal complexes; and (iii) their applications to organic/organometallic transformations.
Related to item (ii), there are some reports in the literature concerning the development of complexes containing two phenolate groups and two nitrogen atoms. 15s exemplified by the work of Strautmann et al., 16 the main goal in this area is to understand the electrochemical and spectroscopic features of these complexes.It has been shown that the redox processes observed in positive potential are ligand centered, resulting in phenoxyl radical species.Relevant to this subject, iron complexes containing tertiary diamines connected to 3,5-bis-tert-butylphenol have been characterized by spectroscopic, electrochemical and structural methods. 6,16,17Among these iron complexes, just one of them contains a diazocyclic backbone (1,4-bis(2hydroxy-3,5-di-tert-butylbenzyl)-1,4-diazepane). 17The others were modeled on ethanediamine and o-phenylene diamine. 6,16ue to the restricted number of iron complexes containing diazocyclic units as central core and phenolate groups as pendant arms in the ligand structure, this work was designed to investigate the physical-chemical and structural features of iron complexes containing different diazocycle-bis(di-tert-butyl-phenol) ligands.Thus a family of diazocyclic ligand containing piperazine (H 2 L1), homopiperazine/1,4-diazepane (H 2 L2), hexahydropyrimidine (H 2 L3) and hexahydropyrimidin-5-ol (H 3 L4) (Figure 1) was prepared and the spectral (infrared (IR), UV-Vis, electron paramagnetic resonance (EPR)) and electrochemical characterization of their iron complexes were carried out.Furthermore, the molecular structure of the iron complexes containing H 2 L2 and H 3 L4 are presented.The radical species formed were confirmed by coulometry and monitored by EPR and UV-Vis (in situ) spectroscopies.In this study we have observed that the storage of oxidizing equivalents on the ligand was easiest in complexes with lower Lewis acidity.

Materials and general methods
All reagents and solvents for syntheses and analyses were of analytical and/or spectroscopic grade and used without further purification.Elemental (C, H, N) analysis was performed on a Thermo Scientific Flash 2000 CHN analyzer.Melting points were performed on a Fisatom melting point apparatus, model 430.Infrared spectra were recorded in KBr disks on a Shimadzu IRAffinity-1 and the electronic spectra (200-1100 nm range) were recorded in CH 2 Cl 2 with a UV-Vis Cary 50 Bio Varian.The nuclear magnetic resonance (NMR) analyses were carried out in a Jeol Eclipse 400+ operating at 400 MHz for 1 H and 100 MHz for 13 C. Cyclic voltammograms (CVs) were carried out with an Autolab PGSTAT 10 potentiostat/galvanostat in dichlorometane containing 0.1 mol L -1 tetrabutylammonium perchlorate (TBAClO 4 ) as supporting electrolyte under an argon atmosphere at room temperature.The electrochemical cell employed was a standard three-electrode configuration with a platinum-wire as auxiliary electrode, a platinum wire as pseudo-reference electrode and a glassy carbon disk as the working electrode, for the complexes [FeL1]Cl (1), [FeL3]Cl (3) and [FeL4] (5).For complexes [FeL2Cl]  (2) and [Fe 2 (HL4)L4Cl] (4) a platinum disk was employed as working electrode.The ferrocenium/ferrocene redox couple was used as internal standard (0.400 V vs. normal hydrogen electrode (NHE)). 18The potentials were corrected and are given vs. Fc/Fc + .The electrochemical formation of the phenoxyl radical species (bulk electrolysis) was carried out by using an Autolab PGSTAT 10 potentiostat/ galvanostat in dichlorometane containing 0.1 mol L -1 tetrabutylammonium perchlorate (TBAClO 4 ) as supporting electrolyte under an argon atmosphere at 25 o C, and a three electrode system.The working electrode was a reticulated vitreous carbon electrode; the counter electrode and the pseudo-reference electrodes were platinum wires.Before the coulometry, cyclic voltammetry experiments were carried out to determine the potential range in which the ligand oxidation was taking place.Different potential were applied and spectral changes in the visible range were followed by in situ measurements employing a fiber optic probe connected to the Varian Cary 50 UV-Vis spectrophotometer.Samples of the solutions were frozen (liquid nitrogen) immediately before and after the electrochemical experiments and their EPR spectra were obtained at X-band frequency (9 GHz), at 100 K, using a Bruker E500 spectrometer with a high sensitive cylindrical cavity, and the following conditions: microwave power of 5 mW; modulation frequency of 100 kHz; modulation amplitude of 1 G.The Qpow program was used to simulate EPR spectra. 19The g-factor was referenced by a MgO:Cr 3+ (g = 1.9797) sample attached to the sample to be analyzed.The electrical conductivity of the complexes was measured with a Biocristal conductivity meter, in dichloromethane, employing a concentration of 1 × 10 -3 mol L -1 .
The crystallographic analyses for complexes 2 and 4 were carried out with a Bruker GADDS diffractometer with graphite-monochromated Cu-K α radiation at 110.15 K, using Olex2. 20The structure was solved with the ShelXS 21 structure solution program using direct methods and refined with the ShelXL 22 refinement package using least squares minimization.All non-hydrogen atoms were refined anisotropically.H atoms attached to C atoms were placed at their idealized position, with C-H distances and U eq values taken from the default settings of the refinement program.The H atom from the phenol group presented in 4 was found in the Fourier difference map and treated with a riding model.One of the tert-butyl groups observed in 4 was found to be disordered and was modeled with rigid bond restraints, with σ for 1-2 distances of 0.005 and σ for 1-3 distances of 0.005.A dark purple crystal of complex 2 with parallelogram form was isolated from a crystalline sample, while the crystals of complex 4 showed a dark red color, with needle shape.Both crystals were isolated from the reaction media after resting for a few days.Selected crystal and refinement data for both compounds are presented in Table 1, while selected bond lengths and angles are shown in Tables S1 and S2 as supplementary information.
Complex [FeL1]Cl, 1: To a solution of H 2 L1 (0.26 g, 0.5 mmol, 20 cm 3 of dichloromethane) was added a solution of FeCl 3 .6H 2 O (0.24 g, 0.88 mmol, 15 cm 3 of methanol).The resulting brown solution was stirred for 10 min at room temperature.The solution was filtered, concentrated, and a new aliquot of dichloromethane was added.The purple solution was filtered and concentrated again, resulting in a purple solid.Yield: 0. To a solution of H 3 L4 (0.54 g, 1 mmol, 30 cm 3 of ethyl acetate), 0.27 g (1 mmol) of FeCl 3 .6H 2 O in 20 cm 3 of methanol was added, followed by the addition of triethylamine (0.81 g, 8 mmol), resulting in a red solution which was stirred for 1 h at room temperature.After 1 day, a needle-shaped red microcrystalline solid was formed and isolated by filtration.Yield: 0.

Synthesis and general characterization
][32][33][34] An added goal of this study is to evaluate the influence of the ring size, as well as the position of the nitrogen atoms, in the diazocyclic unit on the physical-chemical and structural properties of the iron(III) center.Thus, the ligands H 2 L1, H 2 L3 and H 3 L4 possess a six-membered diazocycle ring, whose difference is in the position of the nitrogen atom in the cycle: positions 1,4 for H 2 L1, 1,3 for H 2 L3 and H 3 L4.The two latter are distinct due to the presence, in H 3 L4, of an alcohol group bound to the central carbon atom located in the propanediamine backbone.On the other hand, the ligand H 2 L2 contains a seven-member ring, whose nitrogen atoms are at the 1,4 positions.
From the four ligands employed in this study, three of those were previously published in the literature. 17,23-25H 3 L4 is a new ligand and it was obtained using the same methodology employed in the synthesis of H 2 L3.Ligands H 2 L2, H 2 L3 and H 3 L4 formed iron complexes that are very stable in the solid state, as well as in solution.On the other hand, the isolation of the complex containing the ligand H 2 L1 was very difficult.Its complex, 1, is very unstable in protic solvent, as was observed by the fast loss of its deep purple color in alcohol solution.The same behavior is observed in CH 3 CN, DMSO, ethyl acetate and acetone.A longer stability was achieved in CH 2 Cl 2 .Although there is in the literature some iron complexes containing a six membered diazocyclic unit (piperazine), 35 this work reveals that six-membered 1,4-diazocyclic ligand containing phenol groups as pendant arms are less stable than those containing nitrogen atoms at the 1,3 positions (H 2 L3 and H 3 L4) or than the seven-membered diazocyclic backbone (H 2 L2).The synthesis, UV-Vis and electrochemical characterization of 2 were described previously, 17 but no structural data was available at that time.
When the similar ligand 1,4-bis-(2-hydroxybenzyl)-1,4-diazepane (a ligand similar to H 2 L2, but without the tert-butyl groups) reacted with FeCl 3 , a dinuclear m-oxo complex was obtained and characterized by X-ray diffraction. 17The metal-ligand bond distances are similar to those observed in 2. The presence of tert-butyl groups is most likely responsible for the lack of dimer formation in this instance.
Compound 4 shows an unsymmetrical dinuclear structure (Figure 4).It contains two molecules of the ligand H 3 L4, two iron(III) centers and one chloro ligand.The iron(III) ions (Fe1 … Fe2 = 3.195 Å) are connected through two alkoxide bridges (O3, O4), which come from the ligands.The metal centers present different coordination arrangements.Fe1 is six-coordinate while Fe2 is fivecoordinate.The ligand molecule around Fe1 is coordinated by all of its donor atoms (N 2 O 3 ).In contrast, the ligand bonded to Fe2 displays a NO 2 coordination set, leaving an amine (N4) and a phenol (O6) groups away from the iron center.There is a chloro ligand completing the Fe2 coordination sphere.

Infrared and electronic spectroscopies
The infrared spectra of the ligands are very similar.All of them display the bands typical of aromatic ring (1450-1480 cm -1 ), the C-O stretching characteristic of phenol group (1235 cm -1 ) and a broad band centered about 3000 cm -1 , attributed to O-H group making hydrogen bonds.In this same range, it is possible to observe some peaks attributed to the methyl and methylene groups.Upon complexation, the broad band disappears, with narrow bands remaining, which confirm the presence of tert-butyl groups and methylene units as well.Above 3200 cm -1 , compounds 1 and 3 present a broad band typical of O-H stretching, which may be due to the presence of solvent molecules as observed in the elemental analysis.On the other hand, this region has no bands for compounds 2 and 5, while for compound 4, there is a very narrow strong band, which is typical of O-H group lacking hydrogen bonding, as that observed in the X-ray molecular structure.The position of the C-O stretching vibration does not change significantly after complexation.The electronic spectra were obtained in dichloromethane.The spectra of the ligands show similar features: a broad band centered at 283 nm that is clearly composed of two overlapping bands whose maxima can be observed near 280 and 285 nm, which are attributed to intraligand p → p* transitions.The e values are close to 4.6 × 10 3 L mol -1 cm -1 for the ligands H 2 L1, H 2 L2 and H 3 L4, and 6.3 × 10 3 L mol -1 cm -1 for H 2 L3.The spectra of 1 and 3 show four bands, whereas the spectra of 2, 4 and 5 present 3 bands between 250-1100 nm.The most intense bands observed are between 240-320 nm, which are typical of intraligand charge transfer p → p* (Table 2).The bands observed in the range 330-360 nm are assignable as ligand to metal charge transfer (LMCT) phenolate pp → Fe III dσ*.The less intense bands, which are observed in the range 480-550 nm, are assignable to (LMCT) phenolate pp → Fe III dp*. 6,7,16,17This transition accounts for the purple color observed for 1, 2 and 3, the pink color for 4 and the red one for 5.It is intriguing that compounds 1 and 3 show a different number of bands (four bands) in the UV-Vis spectra when compared with compounds 2, 4, and 5 (three bands).In the electronic characterization of the compounds, FeLCl and FeLNO 3, 6 which exhibit five and six-coordination geometries, it was also observed only 3 bands in dichloromethane solution.Theoretical calculations of these compounds (FeLCl and FeLNO 3 ) supported that the two bands of lower energies are LMCT, and the bands at higher energy are intraligand transitions. 6Thus, the difference in the number of bands observed in the electronic spectra of the compounds suggests that the coordination environment and the geometry for 1 and 3 should be distinct from 2, 4 and 5. Another interesting point is that the e value for the band at lowest energy observed for 1, 3, and 5 are much lower (< 1000 L mol -1 cm -1 ) when compared with the band observed for 2 and 4 and for the compound FeLCl (5400 L mol -1 cm -1 ). 6This suggests that the overlap between the phenolate pp and iron(III) dp are very poor for the formers, probably due to steric hindrance.On the other hand, we can assume that the overlap involving the orbital phenolate pp and iron(III) dσ are very effective, as the e values of the bands that involve these orbitals demonstrate.

Electrochemistry
Cyclic voltammetry and conductimetry were employed in the electrochemical characterization of the complexes (Table 2 and Figure 5).Conductimetric analyses showed that complexes 2, 4 and 5 are neutral species in CH 2 Cl 2 solution, while compound 3 shows a conductivity measure in the range of 1:1 electrolyte type. 42Complex 1 showed a value between neutral and 1:1 species, but, based on the cyclic voltammetry data (see below) we are proposing that it is a 1:1 species.
The cyclic voltammetry of all the complexes showed redox process attributed to the metal center (Figure 5, processes E 1 and E 1 ').The voltammograms for complexes 1, 2, 3 and 5 are typical of mononuclear complexes, since they show only one redox metal-centered process in the range of −2.0 to 0.3 V.This is in agreement with other analyses, including the X-ray diffraction data for complex 2, which also confirms the formation of mononuclear complex.On the other hand, the cyclic voltammetry for complex 4 shows two redox processes (E 1 and E 1 ') at negative potential range, which is characteristic of a dinuclear complex and is in accordance with the X-ray molecular structure.The E 1/2 for the mononuclear complexes decreases in the order 1 ≈ 3 > 2 > 5, which means that complexes 1 and 3 show the most acidic iron(III) centers in this series.Complex 2 presents a redox potential that is about 780 mV more negative than that from complex 1.Since the conductivimetric study for complex 1 suggests that the chloride ion is acting as a counter ion in 1, the difference in the redox potential between these two complexes may be ascribed to the presence of the chloro ligand bonded to the iron ion in 2 (as confirmed by X-ray analysis) and its absence in 1, since the coordination of the chloro ligand decreases the Lewis acidity of the iron center.This also supports the fact that compound 1 show the most acidic iron center in this series.Interestingly, it was observed that the iron center in 3 shows a redox potential very similar to that observed in 1, which suggests that the coordination environment of the iron centers in both complexes should be similar (N 2 O 2 ), and supports the claim that these complexes should be 1:1 electrolyte type, where the chloride species is acting as counterion and not as ligand.
On the other hand, the voltammogram observed for complex 4 indicates that the coordination environment around the iron centers is kept in the mixed valence as well as in the totally reduced species, since both redox couples (Fe III 2 /Fe III Fe II and Fe III Fe II /Fe II 2 ) showed a quasi-reversible electrochemical behavior.Complex 5 presents the least acid iron(III) center, reflecting the presence of the alcohol group in the ligand structure and its coordination as alkoxide to the metal ion.
When the electrochemical analyses were conducted above 0.5 V, redox processes (E 2 , E 3 ) attributed to the oxidation of the ligand (di-tert-butylphenolate) were observed for all the complexes.]16 Compounds 1 and 3 showed just one irreversible oxidative process at 783 and 998 mV, respectively.Above these potential we have observed a constant increase of the current (data not shown) without any other defined oxidation wave.Wieghardt and co-workers have proposed that the lack of reversibility of these processes indicates that a chemical reaction (polymerization) can be induced by the electrochemical process, resulting in the extinction of the radical species. 43This similar behavior also supports that the coordination environment around the metal center in 1 and 3 should be similar as mentioned above.On the other hand, compound 2 showed two quasi-reversible oxidative processes centered on the ligand, indicating the formation of the species [FeL2Cl] + and [FeL2Cl] 2+ .The dinuclear compound 4 showed two close oxidation waves and just one broad reduction process at positive range.It is possible that this reductive wave is composed of two electron reduction.Compound 5 exhibited two well-defined redox processes and their redox potential are the lowest of the set of compounds described in this study.These observations may be explained by the fact that 5 represents the least acidic iron center and that, due to this, the electron density is preferentially localized on the aromatic rings which makes easier to carry out their oxidation.

Radical species formation and characterization
Due to the redox processes observed at positive potential, bulk electrolysis were carried out aiming to confirm if at these potential, a radical species would be formed.The experiments were followed by in situ UV-Vis spectroscopy as well as by EPR. Figure 6 shows the spectral changes in the UV-Vis spectra of the complexes 2, 4 and 5 during the bulk electrolysis.Only for complex 2 it was possible to follow a significant change in the spectrum during electrolysis.The band at 552 nm decreased and a new band near 420 nm appeared, which is typical of phenoxyl radical 6,7,16 as described previously.For complexes 4 and 5, the spectra do not change significantly, but it was possible to observe a small tendency to form a new band close to 420 nm.On the other hand, no change was observed in the spectra of complexes 1 and 3 after electrolysis.It is important to point out that usually this phenoxyl radical UV-Vis absorption is clearly detected at low temperature 6,16 due to the high reactivity of these species, and this may explain why it was not possible to observe clearly this band for compounds 1, 3, 4 and 5.However, for compound 2, this spectral signature could be characterized at 25 o C, indicating the high stability of the radical species formed.
In an attempt to confirm the formation of phenoxyl radical species in compounds 1, 3, 4 and 5, EPR spectra (100 K) were taken before and immediately after electrolysis (Figure 7).In all spectra, we observed lines characteristic of iron(III) in octahedral site at g = 2, and in orthorhombic site at g = 4.After electrolysis, an intense and isotropic signal, with Gaussian line shape and linewidth 1.3 mT, is observed at g = 2.0045 for complexes 2, 4 and 5.The parameters of this line is consistent with the presence of phenoxyl radical species. 6,16s observed in in situ UV-Vis experiments, EPR spectroscopy also could not detect the presence of radical species for complexes 1 and 3 after electrolysis, although the cyclic voltammetry study had shown that these complexes present related electrochemical processes involving ligand oxidation.This lack of detectable radical signature indicates that the radical species formed after

Conclusions
This work clearly reveals that the diazocyclic units employed in this study were able to coordinate to iron(III) salts, resulting in complexes with free coordination site.However, complex 1, which was synthesized with the piperazine backbone (1,4-diazocyclohexane) ligand shows lower stability when compared with others whose nitrogen atoms are at 1,3 position (H 2 L3 and H 3 L4) or 1,4-diazocycloheptane.Of interest also is the fact that the chemical analyses indicate that both ligand H 2 L1 and H 2 L3, which have similar ring sizes (six-membered), but with the nitrogen atoms at different positions (1,4 vs. 1,3), form iron compounds with similar spectroscopic and electrochemical features, implying similar structures of the iron complexes.Our results indicate that the presence of the alcohol group in H 3 L4 is of importance when designing ligands containing redox active groups, since its coordination as alkoxide decreased the Lewis acidity of the metal center and aids in the oxidation of the phenolate groups as observed by the electrochemical potential presented by complex 5.However, the complex which shows the most stable radical species is complex 2, whose ligand contains a 7-member diazocycle ring and whose phenoxyl radical could be detected at 25 o C by UV-Vis spectroscopy.Furthermore, our data shows that the electrochemical potential to store one oxidizing equivalent on the ligand's structure may be tuned by changing the diazocycle ring and by the presence of good donor groups on the ligand structure.Studies are being carried out at the moment to evaluate if this same behavior is observed with other metal compounds and will be the subject of a future publication.

Figure 1 .
Figure 1.Diazocycle units employed in the synthesis of the ligands.

Figure 3 .
Figure 3.A perspective view of molecular structure of 2 with atomic labeling.The ellipsoids are shown with 50% probability level.Hydrogen atoms have been omitted for clarity.

Figure 4 .
Figure 4.A perspective view of molecular structure of 4 with atomic labeling.The ellipsoids are shown with 50% probability level.Hydrogen atoms and tert-butyl groups have been omitted for clarity.

compounds 1 and 3
oxidation undergo a fast chemical reaction, resulting in the extinction of phenoxyl radical.

Figure 7 .
Figure 7. EPR spectra of the complexes 1 to 5 measured in CH 2 Cl 2 (100 K) before (left) and after (right) electrolysis.Some spectra before electrolysis were enlarged by the factor shown in parentheses.

Table 2 .
UV-Vis and electrochemical a data for complexes 1