Iron(II) Complexes with Scorpiand-Like Macrocyclic Polyamines

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Iron(II) Complexes with Scorpiand-Like Macrocyclic Polyamines: Kinetico-Mechanistic Aspects of Complex Formation and Oxidative Dehydrogenation of Coordinated Amines M. Paz Clares,*,† Laura Acosta-Rueda,‡ Carmen E. Castillo,‡ Salvador Blasco,† Hermas R. Jiménez,§ Enrique García-España,*,† and Manuel G. Basallote*,‡ †

Departamento de Química Inorgánica, Universidad de Valencia, Instituto de Ciencia Molecular, Edificio de Institutos de Paterna, C/Catedrático José Beltrán 2, 46980 Paterna, Valencia, Spain ‡ Departamento de Ciencia de los Materiales e Ingeniería Metalúrgica y Química Inorgánica, Instituto de Biomoléculas (INBIO), Facultad de Ciencias, Universidad de Cádiz, Avda República Saharahui s/n, Puerto Real, 11510 Cádiz, Spain § Departamento de Química Inorgánica, Facultad de Química, Universidad de Valencia, C/Dr. Moliner 50, 46100 Burjasot, Valencia, Spain S Supporting Information *

ABSTRACT: The Fe(II) coordination chemistry of a pyridinophane tren-derived scorpiand type ligand containing a pyridine ring in the pendant arm is explored by potentiometry, X-ray, NMR, and kinetics methods. Equilibrium studies in water show the formation of a stable [FeL]2+ complex that converts to monoprotonated and monohydroxylated species when the pH is changed. A [Fe(H−2L)]2+ complex containing an hexacoordinated dehydrogenated ligand has been isolated, and its crystal structure shows the formation of an imine bond involving the aliphatic nitrogen of the pendant arm. This complex is low spin Fe(II) both in the solid state and in solution, as revealed by the Fe−N bond lengths and by the NMR spectra, respectively. The formation rate of [Fe(H−2L)]2+ in aqueous solutions containing Fe2+ and L (1:1 molar ratio) is strongly dependent on the pH, the process being completed in times that range from months in acid solutions to hours in basic conditions. However, detailed kinetic studies show that those differences are caused, at least in part, by the effect of pH on the rate of formation of the unoxidized [FeL]2+ complex. In this sense, the protonation of the donor atoms in the pendant arm of the scorpiand ligand leads to the formation of protonated species resistant to oxidative dehydrogenation. Complementary studies in acetonitrile solution indicate that the initial stage in the oxidative dehydrogenation process is the oxidation of the starting complex to form a [FeL]3+ complex, which then undergoes disproportionation into [Fe(H−2L)]2+ and [FeL]2+. Experiments starting with Fe(III) have allowed us to determine that disproportionation occurs with first order kinetics both in water and acetonitrile solutions. However, whereas a significant acceleration is observed in water when the pH is increased, no effect of the addition of acid or base on the rate of disproportionation is observed in acetonitrile. Oxidative dehydrogenation of the Fe(II) complex formed in experiments starting with an Fe(III) salt is slower than that occurring when an Fe(II) salt is used, an observation that can be explained in terms of the formation of two different Fe(III) complexes, one of them with a structure unable to evolve directly toward the product of oxidative dehydrogenation.

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INTRODUCTION Macrocyclic polyamines containing a pendant arm with donor groups are often called scorpiand ligands. This term was coined by L. Fabbrizzi to describe the movement of the pendant arm toward the macrocycle core to coordinate the metal ion following a pH change.1 We have observed that similar movements can occur also in the free ligands driven by hydrogen bond formation and π-stacking interactions.2−4 In addition, in recent years some reports have been made showing that several of these ligands and their metal complexes have biological activity which may depend on the conformation of the ligand.5−12 Given those precedents, we decided to explore © XXXX American Chemical Society

the complexation of these ligands with iron, the most biologically abundant transition metal ion. During the course of the work with the scorpiand ligand L,2 (Scheme 1) an Fe(II) complex containing the dehydrogenated ligand H−2L was isolated. The complex, which is the result of a single oxidative dehydrogenation process, is quite stable in aqueous solution. These features led us to carry out an in-depth study aimed at obtaining information about the kinetics and mechanism of Received: December 15, 2016

A

DOI: 10.1021/acs.inorgchem.6b03070 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry

The present Article includes synthetic, structural, thermodynamic, and kinetic data that provide information on both the formation of the Fe(II)-L complex and the subsequent oxidative dehydrogenation. The results indicate that the scorpiand-type nature of the ligand introduces peculiarities associated with the role of the pendant arm in both processes.

Scheme 1. L and H−2L Ligands

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EXPERIMENTAL SECTION

L was synthesized as described in ref 2, and it was isolated as its hydrochloride salt. The Fe(II) and Fe(III) sources were the perchlorate salts for aqueous solutions and the triflate salts for experiments in acetonitrile; all of them were obtained from Aldrich. Other reagents were from different commercial sources and used as received. Caution! Perchlorate salts can be explosive and have to be handled with care. NMR Measurements. The 1H NMR spectra were recorded on Bruker DPX-300 or Bruker 400 AV spectrometers operating at 299.95 MHz or 399.91 for 1H, the solvent signal being used for reference. Adjustments to the desired pH were made using drops of DCl or NaOD solutions. The pD was calculated from the measured pH values using the correlation, pH = pD − 0.4.42 Equilibrium Studies. The potentiometric titrations were carried out at 298.1 ± 0.1 K using 0.15 M NaCl as the supporting electrolyte. The experimental procedure (buret, potentiometer, cell, stirrer, microcomputer, etc.) has been fully described elsewhere.43 The reference electrode was an Ag/AgCl electrode in saturated KCl solution. The glass electrode was calibrated as a hydrogen-ion concentration probe by titration of previously standardized amounts of HCl with CO2-free NaOH solutions and the equivalent point determined by Gran’s method,44−48 which gives the standard potential, E°′, and the ionic product of water (pKw= 13.73(1)). The acquisition of the emf data was performed with the computer program PASAT.46 The computer program HYPERQUAD was used to calculate the protonation and stability constants.47,48 The pH range investigated was 2.5−11.0, and the concentration of Fe(II) and of the ligand ranged from 1 × 10−3 to 5 × 10−3 M with Fe(II)/L molar ratios varying from 2:1 to 1:2. The different titration curves (at least two) were treated either as a single set or as separated curves without significant variations in the values of the stability constants. Finally, the sets of data were merged together and treated simultaneously to give the final stability constants. Crystallographic Analysis. An analysis of single crystals of [Fe(H−2L)](ClO4)2 was carried out with an Enraf-Nonius KAPPA CCD single-crystal diffractometer (λ = 0.71073 Å). The structure was solved using the program SHELXS-86.49 Structural refinement was performed by means of the program SHELXL-97.50 Molecular plots were produced with the program MERCURY51 or ORTEP.52 Crystal data, data collection parameters, and results of the analysis are listed in Table S1. Further details may be obtained from the Cambridge Crystallographic Data Center on quoting the depository number CCDC 1521794. Copies of this information may be obtained free of charge from http://www.ccdc.cam.ac.uk. Kinetic Studies. Kinetic studies were carried out by monitoring the UV−vis spectral changes with time with either an Applied Photophysics SX-17MV stopped-flow instrument provided with a PDA-1 diode-array detector or a Cary 50 Bio conventional spectrometer, depending on the time scale of the reaction. The experimental conditions (solvent, temperature, concentrations, and supporting electrolyte) are given in the Experimental Section for each particular case. The analysis of the kinetic data was carried out with the SPECFIT program,53 which provides the values of the rate constants for each step as well as the calculated spectra for the reaction intermediates in cases of polyphasic kinetics. Electrochemical Studies. Cyclic voltammograms (CV) were recorded with an EG&G Princeton Applied Research model 263A potentiostat/galvanostat using platinum or carbon working electrodes and a platinum wire auxiliary electrode. The reference electrode was Ag/AgCl (saturated KCl) for experiments in water solution, and Ag/

both the coordination of the ligand to the metal center and the oxidative dehydrogenation process. Oxidative dehydrogenation of amines coordinated to metal centers, represented in a simplified way in Scheme 2, has been Scheme 2. Oxidative Dehydrogenation of a Coordinated Amine

comprehensively studied over the last decades. Since the initial report by Curtis on the nickel complex of a tetraazamacrocyclic ligand,13 many examples of oxidative dehydrogenation of amines have been reported with a variety of metals such as Co, Fe, Cu, Rh, and others.14−17 While for most ligands the process leads to an imine complex, oxidative dehydrogenation of monodentate primary amines continues to produce the corresponding nitriles.18−21 Although the studies commonly focus on the process itself and on the nature of the products, in recent years a variety of oxidative dehydrogenation processes catalyzed by metal complexes suitable for synthetic purposes have been also reported.22−25 For those reasons, a deeper understanding of dehydrogenative oxidation processes in metal complexes and a way of controlling them would be desirable. The literature includes examples of dehydrogenative oxidation in iron complexes with a variety of ligands that range from simple open-chain aliphatic amines26 as en to macrocyclic amines as cyclam and related ligands.27−29 Interestingly, there is a significant number of examples of oxidative dehydrogenations including complexes of ligands containing 2-(aminomethyl)pyridyl groups30−34 because of the stability achieved by resonance in the product. Initial oxidation of the metal center is a common mechanistic proposal not only for the case of Fe(II) complexes26,35−38 but also for other metals.39 The oxidized complex can then undergo an OH− assisted deprotonation to form a species that can be considered to contain a ligand-centered radical and that converts to the final dehydrogenated product by transferring one proton and one electron.30−34 The oxidant in the latter process can be also an external oxidant or a second molecule of the intermediate, which undergoes disproportionation to form two Fe(II) complexes; one of them is the product of oxidative dehydrogenation and the other one is the starting complex with the nonoxidized ligand.40,41 B


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Inorganic Chemistry AgNO3 (0.01 M AgNO3 and 0.1 M Bu4NClO4) for experiments in acetonitrile.

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RESULTS AND DISCUSSION Potentiometric Studies on the Formation of Fe(II)-L Complexes. The stability of the Fe(II)-L complexes was determined from potentiometric titrations, and the results are included in Table 1. In addition to the [FeL]2+ species, the Table 1. Stability Constants for the Formation of Fe(II)-L Complexes in Aqueous Solution (25.0 ± 0.1 °C, 0.15 M NaCl) reaction

log K

L + Fe ⇆ FeLa FeL + H ⇆ FeHL FeLH + H ⇆ FeH2L FeL + H2O ⇆ FeL(OH) + H

18.43(2)b 5.04(3) 4.11(2) −10.81(5)

Charges are not shown for simplicity. bThe figures in parentheses show the standard deviation in the last significant figure. a

Figure 2. Ball and stick view of the crystal structure of the complex cation [Fe(H−2L)]2+ showing the coordination environment of the metal ion. The imine bond is formed between atoms N4 and C9.

equilibrium model includes its mono and diprotonated forms, and the monohydroxo [FeL(OH)]+ species. The species distribution curves in Figure 1 show that [FeL]2+ is the

Table 2. Selected Distances (Å) and Angles (deg) for the [Fe(H−2L)]2+ Cation bond distances (Å) Fe1 Fe1 Fe1 Fe1 Fe1 N4

N1 N2 N3 N4 N5 C9

1.915(5) 2.053(4) 2.031(6) 1.862(6) 1.967(5) 1.27(1)

angles N1−Fe1−N2 N1 Fe1 N3 N1 Fe1 N5 N3 Fe1 N2 N4 Fe1 N2 N5 Fe1 N2 N4 Fe1 N3 N4 Fe1 N5

81.41(12) 96.9(2) 96.3(2) 84.82(11) 98.75(12) 97.15(11) 84.6(3)) 82.3(3)

complexes with ligands containing both aliphatic and pyridine donors.30,34,54,55 It is also interesting to note that the complex can be stored for long periods of time without any evidence of subsequent dehydrogenation processes affecting other carbon− nitrogen bonds, in contrast with literature examples showing more extensive dehydrogenation with formation of several imine bonds and even dimeric species resulting from coupling between two macrocyclic units.27−29 The coordination environment of Fe(II) is close to a regular octahedron (Table 2), with the four nitrogen atoms in the macrocycle leaving two adjacent coordination sites in cis which are occupied by the central nitrogen of the tren subunit and the pyridine nitrogen of the pendant arm (N4−C9, 1.27 Å, Figure 2). The general disposition of the six nitrogen donors around the metal center is similar to that found for the [CuL]2+ complex with the nonoxidized L ligand2 and for the [CuL′]2+ complex, where L′ is a related scorpiand with quinoline instead of pyridine in the tail.3,4 In all cases, the macrocyclic pyridine and the two adjacent NH groups define a plane that is completed with coordination of the aliphatic nitrogen in the tail, the remaining coordination sites being occupied by the tertiary amine and the pyridine in the side arm. As expected, the shortest Fe(II)-N distance in the coordination sphere is that with the imine nitrogen (Fe1−N4) followed by those of the pyridine nitrogen atoms (Table 2). The distances with the secondary amine groups of the macrocycle are slightly longer. The 1H NMR spectrum of [Fe(H−2L)]2+ in D2O (Figure S1) is typical of a diamagnetic d6 LS complex, the formation of the

Figure 1. Species distribution curves for the formation of Fe(II)-L complexes in aqueous solution containing equimolar amounts of the metal and the ligand (25.0 ± 0.1 °C, 0.15 M NaCl).

major species in the ca. 6−11 pH range. More basic solutions contain a mixture of this species and [FeL(OH)]+, and acidic solutions contain mixtures with the protonated [Fe(HL)]3+ and [Fe(H2L)]4+ complexes. Nevertheless, because of the occurrence of oxidative dehydrogenation, these results must be taken with care, especially those in basic conditions. In a previous study, it was found that the stability of the related [CuL]2+ complex is higher (log β = 22.6),2 and an X-ray determination of the crystal structure revealed hexadentate coordination of the ligand in [CuL]2+. In view of the high stability of the [FeL]2+ complex and the crystal structures previously reported for the Cu(II) and Mn(II) complexes of L, a similar hexacoordination can be assumed in the case of Fe(II).2,6 Synthesis and Characterization of the [Fe(H−2L)]2+ Complex. Crystals of [Fe(H−2L)](ClO4)2 were isolated by slow evaporation of a 1 × 10−3 M solution of Fe(ClO4)2·6H2O and L in 1:1 molar ratio at pH ca. 8. The crystal structure shows the occurrence of a dehydrogenative oxidation of the ligand with formation of an imine bond between the aliphatic nitrogen in the side arm and the adjacent carbon close to the pyridine ring (Figure 2). Relevant bond distances are included in Table 2, which show that the length of the bond between N4 and C9 is similar to that found in other Fe(II)-imine complexes.32,34 The Fe−N bond lengths are typical of low spin (LS) Fe(II) C


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Inorganic Chemistry

intermediate with a band centered at 400−446 nm (Figures 5 and 6). As will be shown with more detail below, the rate of

imine bond being revealed by the presence of a signal at 10.06 ppm. On the other hand, the electronic spectrum of [Fe(H−2L)]2+ in water shows two intense bands centered at 370−400 and 578 nm (Figure 3). The higher energy band can

Figure 3. Electronic spectrum of the [Fe(H−2L)]2+ complex in 0.15 M NaCl aqueous solution.

Figure 5. Spectral changes with time for a solution containing equimolar amounts of Fe(II) and L (c0 = 5.0 × 10−5 M) in water solution at pH 8.0 (25.0 °C, 0.15 M NaCl). The changes in the Figure correspond to a time of 8000 min.

be assigned to a Fe(II) to pyridine charge transfer, whereas the band close to 600 nm corresponds to charge transfer from Fe(II) to the new π system involving the imine bond, as observed for related complexes.30,31 Kinetic Studies on the Formation of [Fe(H−2L)]2+: Complex Formation and Oxidative Dehydrogenation. A detailed study of the behavior of solutions containing Fe(II) and L in 1:1 molar ratio was carried out with the purpose of obtaining information about the kinetics and mechanism of the oxidative dehydrogenation process. The spectrum of freshly prepared aqueous solutions containing an equimolar mixture of Fe(II) and L is featureless in the visible region, but it evolves with time to yield the spectrum of the [Fe(H−2L)]2+ complex. Preliminary kinetic studies monitoring the spectral changes under different conditions showed a variety of behaviors. Thus, experiments at pH 2−4 showed the direct formation of [Fe(H−2L)]2+ in a very slow process that lasts for many days (Figure 4). The spectral changes obtained under these conditions can be fitted satisfactorily to a single kinetic step, but the process is so slow that reliable values of the rate constants could not be obtained. In contrast, experiments in neutral and basic solutions showed biphasic kinetics, the formation of [Fe(H−2L)]2+ occurring through a reaction

Figure 6. Calculated spectra from the fit of spectral changes in Figure 5 to a kinetic model with two consecutive exponentials (A → B → C): A (dotted), B (solid), and C (dashed).

formation of this intermediate (k1obs) is independent of the presence of oxygen, but it is very dependent on the pH. The increase of k1obs with pH is so large that in basic solution the rate of the process is at the limit of the stopped-flow technique (typical spectral changes are shown in Figure S2). Nature of the Reaction Intermediate. Regarding the nature of the intermediate, one possibility is that it corresponds to the [FeL]2+ complex, which would be formed slowly in acidic and neutral solutions because of the high protonation state of L under those conditions. Another possibility is that it corresponds to an Fe(III) intermediate formed in the oxidative dehydrogenation process, as mechanistic proposals in the literature for oxidative dehydrogenation of Fe(II) complexes assume an initial oxidation to Fe(III). However, this intermediate is also formed in experiments carried out inside a glovebox under nitrogen atmosphere, thus ruling out the possibility of an oxidized species. So, the most reasonable interpretation appears to be considering that the intermediate detected in the reaction of L with Fe(II) is actually the [FeL]2+ complex. Additional evidence in this sense was provided by the 1 H NMR spectrum of a D2O solution containing Fe(II) and L

Figure 4. Spectral changes with time for a solution containing equimolar amounts of Fe(II) and L (c0 = 5.0 × 10−5 M) in water solution at pH 4.1 (25.0 °C, 0.15 M NaCl). The changes in the Figure correspond to a time of 8000 min. D


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Inorganic Chemistry

Figure 7. Cyclic voltammograms (left) and UV−vis spectra (right) recorded at different reaction times for water solutions containing equimolar amounts of L and Fe(II) (1 mM) in 0.15 M NaCl at pH 2.7 (top) and pH 9.3 (bottom) at a scan rate of 100 mV s−1. Although both UV−vis spectra were recorded with a 10 mm path length, the spectra at the bottom of the Figure were measured for a 0.23 mM solution of L and Fe(II) to avoid saturation.

(see Figure S8) using the same experimental setup as that used for Figure 7. To obtain additional information, some experiments were carried out in acetonitrile. When acetonitrile solutions containing Fe(II) and completely deprotonated L in equimolar ratio are prepared inside a glovebox with a nitrogen atmosphere, a rapid color change is observed, and the electronic spectrum (Figure S4, red line) is essentially coincident with that of the intermediate species formed in aqueous solution (Figure 6, solid line). The electrochemistry in acetonitrile (oxidation at +0.09 V vs Ag/AgNO3 and + 0.39 V vs SCE56) was also similar to that observed in aqueous solution for the intermediate, thus reinforcing the identification of the intermediate as [FeL]2+. When the solution is taken out of the inert atmosphere, the color changes gradually to dark purple, the UV−vis spectrum shows the gradual appearance of the bands typical of [Fe(H−2L)]2+ (Figure S4), an imine signal appears at 10.05 ppm in the 1H NMR spectrum. The cyclic voltammogram (Figure S5) shows the appearance of the oxidation signal at 0.33 V vs Ag/AgNO3 (+0.63 V vs SCE56) corresponding to the final product, thus supporting the interpretation of the results in water solution in terms of a

in 1:1 molar ratio with the pD adjusted to 10 in the absence of oxygen, which is typical of a diamagnetic species and shows signals for all the protons of the ligand (see Figure S3), thus confirming that at this point L has not suffered any degradation and that the [FeL]2+ complex formed is low spin. It should be noticed that in addition to the signals for [FeL]2+, the spectrum in Figure S3 shows also minor signals that can correspond to the [FeL(OH)]+ hydroxocomplex and/or uncomplexed L, as the experiment was carried out with a slight deficit of the metal ion. The reaction was monitored with simultaneous recording of the UV−vis spectrum and the cyclic voltammogram in acid and basic solutions (Figure 7). These experiments allowed the identification of the oxidation peaks for both the intermediate (+0.34 V vs Ag/AgCl, +0.30 V vs SCE) and the final dehydrogenated product (+0.64 V vs Ag/AgCl, +0.60 V vs SCE), but no signal for any complex preceding the intermediate was observed, thus supporting the conclusion that the intermediate is an [FeL]2+ complex formed from Fe2+ and the protonated ligand. In addition to those signals, there is a cathodic peak at c.a. −0.6 V vs Ag/AgCl which can be assigned to oxygen reduction on the basis of the CVs recorded for a 0.15 M NaCl aqueous solution under nitrogen and under oxygen E


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Inorganic Chemistry first kinetic step corresponding to the formation of [FeL]2+ and a second step for its conversion to [[Fe(H−2L)]2+. Kinetics of Formation of the [FeL]2+ Complex. Kinetic studies on the formation of [FeL]2+ in basic solutions were carried out with a stopped-flow instrument by mixing a solution containing equimolar amounts of ligand and Fe(II) at pH 4 with a solution of the ligand buffered at the desired pH. The kinetic data (Figure S2) were analyzed taking into consideration that Fe(II) and the ligand are not under pseudo-first order conditions, i.e., A + B → C kinetic model with A and B at similar concentrations. In all cases, the process occurs in a single kinetic step, and the dependence of the second order rate constant k1obs with pH (Figure 8) is typical of processes

hydrolysis of the iron aqua complex. Despite the experimental and fitting errors, the value derived for kL‑Fe is close to that expected for complex formation between Fe(II) and a neutral ligand according to the Eigen-Wilkins mechanism. Thus, if the equilibrium constant for the formation of an outer sphere complex between reactants with charges +2 and 0 is assumed to be close to the value determined using the Fuoss eq (0.16 M−1)57,58 and the rate constant for ligand coordination is assumed to be close to water exchange in the Fe(II)-aqua complex (4.4 × 106 s−1),59 the rate constant for complex formation can be expected to be ∼7 × 105 M−1 s−1, close to the experimental value. HL+ = L + H+;

KHL

H 2L2 + = HL+ + H+;

H3L3 + = H 2L2 + + H+; L + Fe2 + → FeL2 +; k1obs =

(1)

KH2L

(2)

KH3L

(3)

kL − Fe

(4)

KHLKH 2LKH 3LkL − Fe KHLKH 2LKH 3L + KH 2LKH 3L[H +] + KH 3L[H +]2 + [H +]3 (5)

At this point, it should be noted that a previous study on the kinetics of complex formation of Cu(II) with L in water showed that complexation in acidic media occurs through reaction of the metal ion with the H3L3+ form of the ligand, the process occurring at a rate much slower than that expected from the donor atom with a tail approach to complex formation with polydentate ligands.2 However, the experimental conditions in that study were quite different from those in the present work because the pH in the Cu(II) case was maintained in all cases below 5.5, which makes extremely low the concentrations of unprotonated L. Oxidative Dehydrogenation of [FeL]2+ to Form [Fe(H−2L)]2+. As pointed out above, the conversion of [FeL]2+ to [Fe(H−2L)]2+ is slower than the formation of [FeL]2+, and it can be monitored by conventional UV−vis in neutral and basic aqueous solutions (Figure 5). The spectral changes show that the species with two overlapping bands at 412 and 448 nm is converted to the product of oxidative dehydrogenation with a band at 580 nm and two overlapping bands at 378 and 402 nm (see Figure 6). The changes could be satisfactorily fitted to a single first order kinetic step with the values of the rate constant (k2obs) in Table 3. Attempts to fit to a second order dependence were unsatisfactory. The values of k2obs show more modest

Figure 8. Plot of the pH dependence of the observed rate constants for the formation of the [FeL]2+ complex. Most experiments were carried out at 25.0 °C in the presence of 0.15 M NaCl by mixing in a stopped-flow instrument a solution containing an equimolar mixture of Fe(II) and L at pH 2 with a solution buffered at the desired pH, but the data at the lower pH values were obtained with a conventional spectrophotometer. The solid line corresponds to the fit of the data by eq 5. See Figure S7 for a plot of the [H+] dependence of the observed rate constants.

occurring with an acid−base pre-equilibrium with a pKa somewhat larger than 10. Although most of the data in Figure 8 were obtained using N2 saturated solutions, some experiments using solutions saturated with air or O2 indicated that the rate is independent of the presence of dioxygen. Figure 8 also includes some kinetic data at lower pH obtained using a conventional spectrophotometer. In these cases, the kinetics is biphasic, and the first resolved step corresponds to the formation of [FeL]2+. Inspection of the species distribution curves for the ligand (Figure S6) indicates that at high pH there is conversion of [HL]+ to L with a pK = 10.21,2 thus suggesting that the large increase of k1obs results from the formation of [FeL]2+ through the reaction of Fe(II) with L, the contribution of [HL]+ to the rate of complexation being negligible. However, as the equilibrium data indicate that there are also significant amounts of [H2L]2+ and [H3L]3+ under the conditions of the kinetic experiments, these acid−base equilibria were also included in the kinetic model (eqs 1−4). The observed rate constant for this model is given by eq 5, and a fit of the data to this equation with the values of the equilibrium constants fixed at the potentiometrically determined values leads to a value of kL‑Fe = (4.5 ± 0.4) × 106 M−1 s−1. No improvement in the quality of the fit is obtained by adding contributions from the protonated form of the ligand or the [Fe(OH)]+ species formed by

Table 3. Observed Rate Constants for the Conversion of [FeL]2+ to [Fe(H−2L)]2+ in Aqueous Solutions (25.0 °C, 0.15 M NaCl) pH

buffer

conditionsa

104 k2obs (s−1)

10.1 9 8 7 6 10.1 9 8

borate borate MOPS MOPS MES borate borate MOPS

O2 O2 O2 O2 O2 N2 N2 N2

1.6 ± 0.1 1.5 ± 0.2 1.56 ± 0.05 1.38 ± 0.02 0.85 ± 0.01 1.1 ± 0.1 0.8 ± 0.2 0.57 ± 0.03

a

The experiments were carried out in solutions purged during 5 min with O2 or N2.

F


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Inorganic Chemistry changes with pH than k1obs, which makes k1obs become smaller than k2obs in acid solutions, and so, kinetic information about oxidative dehydrogenation cannot be obtained in acid conditions. The decrease of k2obs when the acidity is increased can be explained by the fact that [FeL]2+ exists in equilibrium with [Fe(HL)]3+ and [Fe(H2L)]4+ (see Table 1 and Figure 1), whose concentrations increase at decreasing pH. As oxidative dehydrogenation involves the amine group in the pendant arm, which would not be coordinated in these species, the process is hindered. In that case, the only way to form [Fe(H−2L)]2+ is through the formation of [FeL]2+, whose lower concentrations in acidic solutions lead to slower oxidative dehydrogenation. It is also important to note that [FeL]2+ is so air-sensitive that the reaction occurs even in solutions bubbled with N2 for several minutes, which indicates that these purged solutions are not oxygen free. Nevertheless, the process is significantly slowed down when most of the oxygen is excluded (see Table 3), which suggests that it is kinetically controlled by oxidation to an Fe(III) complex, which would then disproportionate rapidly to give [Fe(H−2L)]2+ and [FeL]2+. As pointed out above, acetonitrile solutions containing equimolar amounts of Fe(II) and L in an inert atmosphere show the rapid formation of [FeL]2+ that then converts slowly to [Fe(H−2L)]2+ upon exposure to air (Figure S4). Whereas the formation of [FeL]2+ is too fast to be measured, the kinetics of oxidative dehydrogenation is slow enough to be monitored and the spectral changes could be also satisfactorily fitted to a single first order kinetic step with a rate constant of (2.80 ± 0.05) × 10−4 s−1. This value is of the same order of magnitude as those observed in water, thus suggesting that changing the solvent has a minor effect on the rate of oxidation. As expected, oxidative dehydrogenation is completely inhibited when the experiments are carried out inside a glovebox with a nitrogen atmosphere. Because the ligand used for the experiments in acetonitrile was in the completely deprotonated form and there is no source of protons in this solvent, [FeL]2+ cannot form the protonated species that exist in aqueous solutions. Moreover, the absence of acid−base equilibria hinders the possibility of discerning if oxidative dehydrogenation in acetonitrile solution requires deprotonation of the coordinated amine. To obtain some information on this point, the kinetics of oxidative dehydrogenation was measured in acetonitrile solutions containing equimolar amounts of Fe(II) and L in the presence of 1−2 equiv of acid (HCOOH) or base (NaOD). The results indicate that addition of acid or base does not cause any change in the UV−vis spectrum of [FeL]2+, thus showing that in acetonitrile there is no formation of [Fe(HL)]3+ or [FeL(OH)]+, at least under the experimental conditions used. Moreover, the addition of acid or base does not cause any change in the rate constant for the formation of [Fe(H−2L)]2+, the values being (2.8 ± 0.1) × 10−4 s−1 in the presence of acid and (2.75 ± 0.05) × 10−4 s−1 in the presence of base. As experiments starting with [FeL]2+ suggest that the formation of the product of oxidative dehydrogenation is kinetically controlled by oxidation to an [FeL]3+ complex both in water and acetonitrile solutions, additional experiments starting with Fe(III) were carried out in an attempt to obtain some information on the oxidative dehydrogenation process. Experiments in acetonitrile solution carried out inside a glovebox using equimolar mixtures of L and Fe(III) show a rapid color change that corresponds to the formation of an [FeL]3+ complex with an intense absorption in the UV (Figure 9 solid line) region, but this process is too fast for obtaining

Figure 9. UV−visible absorption spectra of a 1 mM acetonitrile solution of Fe(III) and L (1:1) in 0.1 M [(Bu4N)ClO4] recorded immediately after mixing Fe3+ and L, [FeL]3+(solid line), after evolution of [FeL]3+ in the absence of O2 (dotted line), and after 14 h in the presence of O2 (dashed line).

reliable kinetic data with the experimental setup used. The electronic spectrum recorded immediately upon mixing evolves with time to yield a spectrum that clearly shows the formation of a mixture of [FeL]2+ and [Fe(H−2L)]2+ (Figure 9 dotted line), a conclusion further supported by the cyclic voltammograms, which show the oxidation signals corresponding to both species. When the solution is taken out of the glovebox and bubbled with O2, a further slow increase of the [Fe(H−2L)]2+ bands is observed (Figure 9, dashed line). Figure 10 illustrates

Figure 10. Kinetic trace recorded at 584 nm for an acetonitrile solution containing Fe(III) and L (1:1 molar ratio, 1 mM, 25.0 °C). The first part of the curve was obtained inside a glovebox with a nitrogen atmosphere, and the second part was recorded after taking the sample out of the glovebox and exposing to air. The discontinuity is caused by the time required to take the sample out of the glovebox.

the time dependence of the spectral changes in these experiments. The changes observed in the absence of air are fully consistent with disproportionation of a [FeL]3+ complex (eq 6), a process that does not require any oxidant, and so it can occur under the inert atmosphere of the glovebox. The subsequent changes observed in the presence of O2 would then correspond to oxidation of the previously formed [FeL]2+ complex (eq 7), the resulting [FeL]3+ undergoing again disproportionation to form additional amounts of [Fe(H−2L)]2+. Actually, the intensity of the [Fe(H−2L)]2+ bands increase upon exposition to O2 to absorbance values that closely approach the double of the intensity of the bands formed immediately upon mixing the Fe(III) salt and L (Figure 10). Kinetic traces such as those in Figure 10, both before and after exposure to air, were satisfactorily fitted by a single exponential, and the values of the rate constants (kd for the G


Article

Inorganic Chemistry disproportionation step and kox for the final formation of the [Fe(H−2L)]2+ complex) are included in Table S2. As observed in experiments starting with Fe(II), no significant changes in the kinetics of reaction were observed in the presence of added acid or base. The values of kox, which are in all cases in the (1.1−3.4) × 10−4 s−1 range, are similar to those obtained in experiments starting with Fe(II), the differences being probably caused by the different degree of saturation with air in the transfer of the sample from the glovebox. More surprising is the observation of a first order dependence for the disproportionation step (kd), but attempts to fit to a second order dependence led to worse results. Moreover, the first order dependence of this step was confirmed by the independence of the kd value when the complex concentration was changed. Thus, although disproportionation involves two [FeL]3+ species, the rate-determining step appears to be an intramolecular transformation of one of them (eq 8), which is followed by rapid reaction with the second [FeL]3+ (eq 9). 2[FeL]3 + → [FeL]2 + + [Fe(H−2L)]2 + + 2H+;

higher than 8. Unfortunately, the limited set of values available hinders additional analysis of this rate constant. The second resolved kinetic step shows first order kinetics with rate constants of the order of 10−3−10−4 s−1, and it leads to the appearance of the bands for the [Fe(H−2L)]2+ product, whose intensity continues increasing in the third kinetic step (Figure S9). Thus, despite the first order kinetics the second resolved step corresponds to disproportionation of [FeL]3+, and as occurs with the data in acetonitrile solution, attempts to fit the data to a second order dependence were unsatisfactory. Although the k2 values in Table S3 have a significant dispersion, they clearly show a tendency to decrease as the pH increases, which suggests the occurrence of a deprotonation before the rate-determining step. Thus, the disproportionation process (eq 6) would occur through a mechanism of the type indicated in eqs 10−12, where one [FeL]3+ undergoes deprotonation followed by rate-determining reorganization of the resulting [Fe(H−1L)]2+ and rapid reaction with a second [FeL]3+. Unfortunately, the dispersion of the data hinders a detailed analysis of the pH dependence of k2, but these observations are in line with previous reports showing deprotonation of a coordinated NH group during oxidative dehydrogenation processes.38

kd (6)

[FeL]2 + + O2 → [FeL]3 + ;

[FeL]3 + → ([FeL]3 + )* ; 3+

[FeL]

kox

k 2d

(7) (8)

[FeL]3 + = [Fe(H−1L)]2 + + H+;

3+

+ ([FeL] )* → [FeL]2 + + [Fe(H−2L)]2 + +

+ 2H ;

fast

K 2a

[Fe(H−1L)]2 + → ([Fe(H−1L)]2 + )* ;

k 2d

(10) (11)

(9)

[FeL]3 + + ([Fe(H−1L)]2 + )* → [FeL]2 +

The spectral changes recorded in aqueous solution for experiments using equimolar amounts of Fe(III) and L (Figure 11) at different pH values were more complicated than those

+ ([Fe(H−2L)]2 + + H +;

fast

(12)

It is interesting to note that the disproportionation step occurs with first order kinetics both in acetonitrile and water, although in the latter solvent it occurs with initial proton dissociation because this process is favored by the protic nature of the solvent. In the literature, there are examples of oxidative dehydrogenations for which the disproportionation occurs with both second40,41,60 and first30,32,35,61,62 order kinetics with respect to the metal complex, and there are even some cases in which the rate constants for all of the individual steps have been resolved.39 Therefore, the rate constants for the different processes involved in disproportionation appear to be delicately balanced for each particular complex. The third resolved kinetic step in the experiments starting with Fe(III) in aqueous solution is also first order with rate constants in the order of 10−6 s−1, and it involves the formation of additional amounts of [Fe(H−2L)]2+. However, the values of k3 in Table S3 are 2 orders of magnitude smaller than those observed for the oxidation of [FeL]2+ in experiments starting with Fe(II), and the total amount of product of oxidative dehydrogenation is in all cases much smaller than that formed when the experiments are started using an Fe(II) salt. In addition, the calculated spectra in Figure S9 show that the amount of [Fe(H−2L)]2+ formed in the second resolved step in experiments with Fe(III) is smaller than 50% of the total amount of this compound formed at the end of the third step. Moreover, the spectral changes observed in the Fe(III) experiments do not show any evidence of the appearance of the band at 400−446 nm expected for the [FeL]2+ complex (compare Figures 5 and 11). These observations strongly suggest that the [FeL]2+ and [FeL]3+ complexes formed in the experiments in water starting with Fe(II) and Fe(III) probably have a different structure. To obtain additional information

Figure 11. Typical spectral changes showing the formation of [Fe(H−2L)]2+ in aqueous solutions containing equimolar amounts of Fe(III) and L (c0 = 1.5 × 10−4 M, 25.0 °C, 0.15 M NaCl). The present changes were recorded during 118 h for an experiment carried out at pH 8.0.

observed in acetonitrile, although they can be accommodated to a quite similar reaction scheme. In general, a satisfactory fit of the changes requires a model with three consecutive kinetic steps (see rate constants in Table S3 and calculated spectra in Figure S9). The first step is first order with respect to both Fe(III) and L, so that it corresponds to the formation of an [FeL]3+ complex; it is signaled by the disappearance of the OH− to Fe(III) charge transfer band at 350 nm and leads to a spectrum which is essentially featureless. The values derived for the second order rate constant are within 0.7−4.6 M−1 s−1. However, this step becomes faster when the pH increases, and actually, no values of the rate constant could be obtained at pH H


Article

Inorganic Chemistry

Figure 12. Possible structures adopted by the Fe(II) or Fe(III) complexes of L. The macrocyclic unit adopts a bent structure leaving two vacant coordination sites in cis that can be occupied by water or by one or the two nitrogen donors in the pendant arm. However, whereas the structure in a allows for coordination of the additional nitrogen donors without reorganization, the structure in b would be an N-based isomer with the tertiary amine inverted so that the pendant arm is anti with respect to the two aqua ligands, thus hindering hexadentate coordination. Note that in addition to the actual structures in the Figure, the uncoordinated nitrogens can become protonated, thus yielding [Fe(HxL)]z+ species with x = 1,2.

observed at −0.12 V. CVs recorded at longer reaction times show a decrease in the intensity of this reduction signal as well as an increase of the [Fe(H−2L)]2+ oxidation signal (Figure S11d and e). The formation of some amount of the [FeL]2+ complex is only revealed by a very small shoulder at ca. 0.35 V. After 14 days, the CV shows, in addition to the oxidation signal for [Fe(H−2L)]2+, a quasi-reversible reduction signal at −0.15 V that clearly demonstrates the formation of an Fe(III)-L species resistant to oxidative dehydrogenation. At this point, it is important to note that the cyclic voltammograms in Figure 7 for experiments starting with Fe(II) also show at long reaction times a weak signal at ca. −0.15 V that would indicate that small amounts of the same Fe(III) complex resistant to oxidative dehydrogenation are also formed in this case. The whole set of results can be interpreted in terms of initial formation of two different Fe(III)-L complexes in the first kinetic step. One of those complexes, with a reduction potential close to −0.05 V would evolve to the same [FeL]3+ complex formed upon oxidation of the [FeL]2+ species formed with Fe(II) and would undergo rapid disproportionation (second kinetic step), and the other one (major reduction potential close to −0.12 V) would be a complex much more stable toward disproportionation. The latter species could react more slowly to form a closely related species with a reduction signal at −0.15 V and another species suitable for undergoing oxidative dehydrogenation, thus being responsible for the formation of more [Fe(H−2L)]2+ in the third kinetic step. In contrast, when the reaction is carried out using Fe(II), the major product would be the [FeL]2+ complex with oxidation potential at 0.34 V, which oxidizes and then disproportionates to give [Fe(H−2L)]2+ and the same [FeL]2+, although smaller amounts of a different complex that oxidizes to yield an Fe(III) complex resistant to disproportionation are also formed. Unfortunately, the absence of absorption bands in the UV− vis spectra other than those of the initial Fe3+ species and [Fe(H−2L)]2+ hinders a more detailed analysis of the changes occurring during the reaction when the experiment is started with Fe(III), although the absence of any new absorption band in the 300−400 nm rules out the formation of a LFeIII−O-

about this point, additional experiments were carried out. Thus, the reaction between Fe(II) and L in aqueous solution at pH 6.0 in the presence of added oxidants was also monitored. The oxidants selected were [FeCp2]+, which is able to oxidize the [FeL]2+ complex but not free Fe2+ and Ce4+, which can oxidize both [FeL]2+ and Fe2+. The spectral changes (Figure S10) monitored in the presence of one equivalent of [FeCp2]+ indicate that the formation of [Fe(H−2L)]2+ occurs without the formation of detectable amounts of [FeL]2+, although the amount of [Fe(H−2L)]2+ is similar to that formed in the absence of [FeCp2]+, which is consistent with oxidation and disproportionation of [FeL]2+ being faster than its formation. On the other hand, Figure S10 shows that the amount of [Fe(H−2L)]2+ formed with added Ce4+ is significantly lower, in agreement with initial oxidation of free Fe2+ before complexation with L. As expected, [FeL]2+ is not detected in these experiments because it is also oxidized by Ce4+. These results can be interpreted by considering that the reaction of Fe3+ and L leads mainly to the formation of an Fe(III) complex resistant to oxidative dehydrogenation, probably because the pyridine in the pendant arm is not coordinated, although that species would be in equilibrium with minor amounts of the same [FeL]3+ complex formed upon oxidation of [FeL]2+. The formation of small amounts of [Fe(H−2L)]2+ would result from disproportionation of this minor Fe(III) species. Unfortunately, the spectral changes do not provide any information about the possible formation of two different Fe(III)-L complexes because they do not show any bands assignable to those species. On the other hand, the reaction between Fe3+ and L was also monitored by recording CVs as a function of time (see Figure S11). Thus, the CV of a solution prepared with equimolar amounts of Fe(III) and L at pH 6.0 after 10 min shows the complete disappearance of the free Fe3+ signal (Figure S11a) and the appearance of an unresolved reduction signal with two shoulders at ca. −0.05 and −0.15 V (Figure S11b). A weak signal for oxidation of some [Fe(H−2L)]2+ also appears, and it becomes more evident in the CV recorded after 55 min (Figure S11c). At this point, the reduction feature at −0.05 V has disappeared, and a single irreversible reduction signal is I


Article

Inorganic Chemistry FeIIIL dimer of the type frequently found for other Fe(III) complexes.63−65 With regard to the nature of the different species formed, although no definitive conclusions can be obtained with the data available, one possibility is that reaction of the protonated forms of L with Fe(II) leads mainly to an [FeL]2+ complex with a structure similar to that found in [Fe(H−2L)]2+, in which the ligand adopts a bent conformation with the four nitrogen donors in the macrocycle leaving two additional coordination sites in cis that allow for the simultaneous coordination of the NH and py groups in the pendant arm (Figure 12a). Oxidative dehydrogenation would be easily achieved from this structure through oxidation and disproportionation. On the other hand, the major species formed in the reaction with Fe(III) could be an N-based isomer with the tertiary amine inverted, so that the pendant arm is anti with respect to the two aqua ligands, thus hindering coordination of the donor groups in the arm (Figure 12b). In addition, smaller but significant amounts of the complex suitable for undergoing oxidative dehydrogenation would be also formed in the case of Fe(III), and formation of minor amounts of the complex with a structure such as that shown in Figure 12b would also occur in the case of Fe(II). The formation of two different compounds in the reaction of a metal ion with L cannot be considered surprising because the complex conformational changes occurring in the protonated forms of the ligand allow for the occurrence of parallel reactions leading to different compounds.2 An alternative possibility for the structure in Figure 12b would consist of a pseudoplanar arrangement of the four donors in the macrocycle so that the additional coordination sites would be in trans, thus hindering the simultaneous coordination of the two donor groups in the arm. However, the crystal structures available for metal complexes with L and other tren-based scorpiands show a bent conformation of the macrocycle,2,3,5 which suggest that this conformation is more stable. Nevertheless, the adoption of the more strained pseudoplanar conformation cannot be completely ruled out, especially if the structure is distorted, and the metal ion is located out of the plane. In any case, the fact that the initially formed Fe(III) complexes have reduction potentials significantly lower than that of [FeL]2+ can be interpreted by considering that at least one of the additional coordination sites are occupied by water, in agreement with the higher affinity of Fe(III) by oxygen donors. Formation of hydroxocomplexes does not appear probable because the UV− vis spectra do not show the typical OH− to Fe3+ charge transfer band, and the lower reduction potentials are in agreement with findings for related pyridine containing ligands showing that hypodentate coordination causes a decrease in the reduction potential.66 In the second resolved kinetic step, the minor Fe(III) complex formed in the first step would evolve to a complex with the six nitrogen donors of L coordinated, which would undergo rapid disproportionation to yield [Fe(H−2L)]2+ and [FeL]2+. Oxidation and disproportionation of the previously formed [FeL]2+ would lead to formation of more [Fe(H−2L)]2+ in the third step. However, the calculated spectra in Figure S9 show that the amount of [Fe(H−2L)]2+ formed in the third step is larger than that formed in the second step, which indicates that the major Fe(III) complex also contributes to the formation of the complex with the oxidized ligand. Nevertheless, in order to undergo oxidative dehydrogenation this compound requires isomerization to allow for coordination of the NH and py groups in the pendant, and the signal at −0.15 V in the CVs recorded at long reaction times would then

correspond to the Fe(III) complex formed in this process. Isomerization is expected to be slow because it involves the breaking and formation of some of the Fe−N bonds. In acetonitrile solution, the absence of protonated species reduces the possibilities of formation of different complexes and surely also allows for faster conformational changes so that the structure of both the Fe(II) and Fe(III) complexes formed appears to be the same found in[Fe(H−2L)]2+.

■

CONCLUSION The chemistry of Fe(II) complexes with the scorpiand ligand L has been studied in aqueous solution both from thermodynamic and kinetic points of view. In addition to the [FeL]2+ complex, several protonated and hydroxylated species are formed. Kinetic studies indicate that complex formation is very slow in acidic and moderately basic solutions, but the rate increases considerably at high pH. The pH dependence of the rate constant can be interpreted in terms of complex formation by reaction of Fe(II) with the completely deprotonated form of the ligand with a rate constant that agrees well with expectations from the Eigen-Wilkins mechanism. The NMR and UV−vis spectra of the [FeL]2+ species suggest that it is a low spin complex. However, in the presence of air there is oxidative dehydrogenation of the ligand with formation of a [Fe(H−2L)]2+ complex which contains a carbon−nitrogen double bond involving the aliphatic nitrogen in the scorpiand side arm. A similar behavior is observed in acetonitrile solution, and kinetic studies indicate that oxidative dehydrogenation is kinetically controlled by oxidation to form an [FeL]3+ complex, which then undergoes disproportionation to form [FeL]2+ and [Fe(H−2L)]2+. Experiments starting with Fe(III) have allowed us to determine that disproportionation occurs with first order kinetics both in water and acetonitrile solutions. However, whereas a significant acceleration is observed in water when the pH is increased, no effect of the addition of acid or base on the rate of disproportionation is observed in acetonitrile. Oxidative dehydrogenation of the Fe(II) complex formed in experiments starting with an Fe(III) salt is slower than that occurring when an Fe(II) salt is used, and the amount of [Fe(H−2L)]2+ formed is lower. These observations can be rationalized in terms of the formation of two different Fe(III) complexes that differ in the conformation adopted by the ligand. One of the complexes would have a conformation similar to that found in the crystal structure of [Fe(H−2L)]2+, thus allowing for direct oxidative dehydrogenation, and the other one requiring previous isomerization. Another relevant observation from the point of view of the scorpiand nature of the ligand is the inhibition of the oxidative dehydrogenation in the protonated [FeHL]3+ and [FeH2L]4+ species as a consequence of the dissociation of the donor atom in the scorpiand tail.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.6b03070. Additional NMR, CV, species distribution curves, and kinetic results (PDF) X-ray crystallographic data for [Fe(H−2L)](ClO4)2 (CIF) J


Article

Inorganic Chemistry

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L.; Sanchez-Moreno, M.; Garcia-Espana, E. Synthetic single and double aza-scorpiand macrocycles acting as inhibitors of the antioxidant enzymes iron superoxide dismutase and trypanothione reductase in Trypanosoma cruzi with promising results in a murine model. RSC Adv. 2014, 4, 65108−65120. (11) Olmo, F.; Marín, C.; Clares, M. P.; Blasco, S.; Albelda, M. T.; Soriano, C.; Gutiérrez-Sánchez, R.; Arrebola-Vargas, F.; García-España, E.; Sánchez-Moreno, M. Scorpiand-like azamacrocycles prevent the chronic establishment of Trypanosoma cruzi in a murine model. Eur. J. Med. Chem. 2013, 70, 189−198. (12) Serena, C.; Calvo, E.; Clares, M. P.; Diaz, M. L.; Chicote, J. U.; Beltrán-Debon, R.; Fontova, R.; Rodriguez, A.; García-España, E.; García-España, A. Significant In Vivo Anti-Inflammatory Activity of Pytren4Q-Mn a Superoxide Dismutase 2 (SOD2) Mimetic ScorpiandLike Mn (II) Complex. PLoS One 2015, 10, e0119102. (13) Curtis, N. F. Nickel(II) complexes of two isomeric cyclic tetraamines and dehydro-derivatives with one to four imine donor groups. Chem. Commun. (London) 1966, 881−883. (14) Amendola, V.; Fabbrizzi, L.; Mundum, E.; Pallavicini, P. Formation of a dicopper(I) helicate by oxidative dehydrogenation of a monomeric copper(II) polyamine complex. Dalton Trans. 2003, 773− 774. (15) Christian, G. J.; Arbuse, A.; Fontrodona, X.; Martinez, M. A.; Llobet, A.; Maseras, F. Oxidative dehydrogenation of an amine group of a macrocyclic ligand in the coordination sphere of a CuII complex. Dalton Trans. 2009, 6013−6020. (16) Lay, P. A.; Sargeson, A. M. Tris(1,2-ethanediamine) complexes of osmium(IV), osmium(III) and osmium(II): oxidative dehydrogenation reactions. Inorg. Chim. Acta 1992, 198−200, 449−460. (17) Zhou, X.; Day, A. I.; Edwards, A. J.; Willis, A. C.; Jackson, W. G. Facile C−H Bond Activation: Synthesis of the N4C Donor Set Pentadentate Ligand 1,4-Bis(2-pyridylmethyl)-1,4-diazacyclononane (dmpdacn) and a Structural Study of Its Alkyl−Cobalt(III) Complex [Co(dmpdacn-C) (OH2)](ClO4)2·H2O and Its Hydroxylated Derivative [Co(dmpdacnOH-O)Cl](ClO4)2·C3H6O. Inorg. Chem. 2005, 44, 452−460. (18) Bernskoetter, W. H.; Brookhart, M. Kinetics and mechanism of iridium-catalyzed dehydrogenation of primary amines to nitriles. Organometallics 2008, 27, 2036−2045. (19) Diamond, S. E.; Tom, G. M.; Taube, H. Ruthenium promoted oxidation of amines. J. Am. Chem. Soc. 1975, 97, 2661−2664. (20) Keene, F. R.; Salmon, D. J.; Meyer, T. J. Oxidation of primary amines bound to Bis(2,2′-Bipyridine)Ruthenium(II). J. Am. Chem. Soc. 1976, 98, 1884−1889. (21) McWhinnie, W. R.; Miller, J. D.; Watts, J. B.; Waddan, D. Y. Hexakisamine-and benzonitrilepentakisamine-complexes of ruthenium(II). The ruthenium induced activation of amines to oxidation. Inorg. Chim. Acta 1973, 7, 461−466. (22) Hammond, C.; Schümperli, M. T.; Hermans, I. Insights into the Oxidative Dehydrogenation of Amines with Nanoparticulate Iridium Oxide. Chem. - Eur. J. 2013, 19, 13193−13198. (23) Mori, K.; Yamaguchi, K.; Mizugaki, T.; Ebitani, K.; Kaneda, K. Catalysis of a hydroxyapatite-bound Ru complex: efficient heterogeneous oxidation of primary amines to nitriles in the presence of molecular oxygen. Chem. Commun. (Cambridge, U. K.) 2001, 461−462. (24) Samec, J. S. M.; Éll, A. H.; Bäckvall, J.-E. Efficient RutheniumCatalyzed Aerobic Oxidation of Amines by Using a Biomimetic Coupled Catalytic System. Chem. - Eur. J. 2005, 11, 2327−2334. (25) Zhu, B.; Lazar, M.; Trewyn, B. G.; Angelici, R. J. Aerobic oxidation of amines to imines catalyzed by bulk gold powder and by alumina-supported gold. J. Catal. 2008, 260, 1−6. (26) Goedken, V. L.; Busch, D. H. Facile promotion of oxidative dehydrogenation by iron ions and the synthesis of new complexes of iron with highly unsaturated tetraaza macrocycles. J. Am. Chem. Soc. 1972, 94, 7355−7363. (27) Mountford, H. S.; Spreer, L. O.; Otvos, J. W.; Calvin, M.; Brewer, K. J.; Richter, M.; Scott, B. Structure and properties of a bimetallic complex with a conjugated bridge between iron atoms in two macrocycles. Inorg. Chem. 1992, 31, 717−718.

AUTHOR INFORMATION

Corresponding Authors

*(M.P.C.) E-mail: [email protected]. *(E.G.-E.) E-mail: [email protected]. *(M.G.B.) E-mail: [email protected]. ORCID

Manuel G. Basallote: 0000-0002-1802-8699 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS Financial support by the Spanish MICINN and MEC and FEDER funds from the European Union are gratefully acknowledged (grants CSD2010-00065, CTQ2015-65707-C22-P, CTQ2013-14892, CTQ2015-71470-REDT, and CTQ2016-78499-C6-1-R), Unidad de Excelencia Mariá de Maeztu MDM-15-0538), and Generalitat Valenciana (PROMETEO II 2015-002). Helpful suggestions by the referees are also gratefully acknowledged.

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