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Article Cite This: Inorg. Chem. XXXX, XXX, XXX−XXX

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Coordination Chemistry of Cu2+ Complexes of Small N‑Alkylated Tetra-azacyclophanes with SOD Activity Á lvaro Martínez-Camarena,†,⊥ Andrea Liberato,‡,⊥ Estefanía Delgado-Pinar,† Andreś G. Algarra,‡ Javier Pitarch-Jarque,† Jose ́ M. Llinares,†,§ M. Á ngeles Mañez,‡ Antonio Domenech-Carbo,́ ∥ Manuel G. Basallote,*,‡ and Enrique García-España*,†

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†

Instituto de Ciencia Molecular, Departamento de Química Inorgánica, Universidad de Valencia, 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 Orgánica, Universidad de Valencia, C/Dr. Moliner s/n, 46100, Burjassot, Valencia, Spain ∥ Departamento de Química Analítica, Universidad de Valencia, C/Dr. Moliner s/n, 46100, Burjassot, Valencia, Spain S Supporting Information *

ABSTRACT: A new tetraaza-pyridinophane macrocycle (L1) Nalkylated with two isopropyl and one methyl groups symmetrically disposed has been prepared and its behavior compared with those of the unsubstituted pyridinophane (L3) and the related compound with three methyl groups (L2). The protonation studies show that, first, a proton binds to the central methylated amine group of L1, while, second protonation leads to a reorganization of the protons that are at this stage attached to the lateral isopropylated amines. The X-ray structure of [HL1]+ agrees with the UV−vis and NMR studies as well as with the results of DFT calculations. The stability of the Cu2+ complexes decreases on increasing the bulkiness of the alkyl substituents of the amine groups. The crystal structures of [CuL1Cl](ClO4) and [CuL1(H2O)](ClO4)2·H2O show square pyramidal coordination geometries with the ligands disposed in a bent L-shaped conformation. Kinetic studies indicate that the rates of both complexation and ligand dissociation decrease with the bulkiness of the substituents, so that the stability changes are surely the results of compensating effects, complex formation dominating over complex dissociation. The pH dependence of the rate constants for complex formation cannot be explained by consideration of rapid pre-equilibria involving the different protonated forms of the ligand, and it has been interpreted in terms of a mechanism involving an acid−base equilibrium for a reaction intermediate. NBT SOD studies show that the Cu2+ complex of the bulkiest L1 ligand is the one having the highest activity (IC50 = 0.26(5) μM, kcat = 13.7 × 106 M−1 s−1) which can be associated with the poorer σ-donor ability of the tertiary amino groups, and the rigidity of the system, caused by the bulky isopropyl groups.

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INTRODUCTION Tetra-azamacrocycles are among the ligands most widely studied in coordination and supramolecular chemistry.1−4 Initial interest in these ligands was related to their capacity for arranging four nitrogen donors in a square-planar fashion, resembling the organization achieved by porphyrin ligands. On the other hand, the introduction of aromatic rings in the ligand stiffens the macrocycle and changes the spatial arrangement of the donor atoms, thus generating open sites to which exogenous ligands or substrates may bind more efficiently. Indeed, coordinative unsaturation is a requisite for many catalytic sites to work. In this respect, the pyridinophane L35 (Scheme 1) and its derivative having a pyridol ring instead of a pyridine one have been reported as potential drugs for treating neuro-degenerative disorders such as Alzheimer disease.6,7 It has been suggested that these ligands might permeate the © XXXX American Chemical Society

blood brain barrier (BBB) capturing copper ions and developing thereby antioxidant protection.7 On the other hand, some of us have recently reported that the iron complex of an L3 derivative with the three aliphatic nitrogens methylated (L2) somehow mimics cytochrome P450 and catalyzes hydroxylation reactions of saturated alkanes with impressive rates.8,9 In parallel to continuing our work on the iron complexes,8−10 we considered of interest to carry out an indepth study about the protonation and copper(II) coordination behavior in water of this class of ligands. Since as it has been extensively documented N-alkylation of polyamine ligands strongly affects their acid−base behavior, coordination Received: June 8, 2018

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DOI: 10.1021/acs.inorgchem.8b01492 Inorg. Chem. XXXX, XXX, XXX−XXX

Article

Inorganic Chemistry

122.87, 60.48, 54.85, 53.16, 51.61, 40.28, 16.65. Anal. Calc. for C18H32N4·3.3HCl·4H2O·0.4(C4H8O2): C 44.2%, H 8.8%, N 10.5%. Found: C 44.4%, H 8.5%, N 10.6%. MS (ESI) m/z 305.1 [M + H]+. EMF Measurements. The potentiometric titrations were carried out at 298.1 ± 0.1 K using 0.15 M NaClO4 as supporting electrolyte. The experimental procedure (buret, potentiometer, cell, stirrer, microcomputer, etc.) has been fully described elsewhere.20 The acquisition of the electromotive force (emf) data was performed with the computer program PASAT.21 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 was determined by the Gran’s method,22,23 which gives the standard potential, E°, and the ionic product of water (pKw = 13.73(1)). The computer program HYPERQUAD was used to calculate the protonation and stability constants.24 The HYSS25 program was used to obtain the distribution diagrams. The pH range investigated was 2.5−11.0. The concentrations of Cu2+ and of the ligands ranged from 0.5 to 5 mM using Cu2+:L molar ratios of 1:1. The titration curves for each system (at least two titrations, ca. 100 experimental points) were treated either as a single set or as separated curves without significant variations in the values of the stability constants. NMR Measurements. The 1H and 13C NMR spectra were recorded on a Bruker Advance DRX 500 spectrometer operating at 500 MHz for 1H and at 100.6 MHz for 13C. For the 13C NMR spectra, dioxane was used as a reference standard (δ 67.4 ppm), and for the 1H spectra, the solvent signal was used. 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 − 0.4.26 Single-Crystal X-ray Diffraction Analyses. Single-crystal X-ray diffraction (SXRD) data for [HL1](ClO4), [CuL1Cl](ClO4), and [CuL1(H2O)](ClO4)2·H2O were collected on an Xcalibur diffractometer (Agilent Technologies, Sapphire 3 CCD detector) using a single wavelength X-ray source with Mo Kα radiation, λ = 0.71073 Å, and 120(1) K in all cases. The selected single crystals were mounted using Paratone-N hydrocarbon oil27 on the top of a loop fixed on a goniometer head and immediately transferred to the diffractometer. Data collection, analytical absorption correction, and data reduction were performed with the Oxford program suite CrysAlisPro.28 An empirical absorption correction was applied using spherical harmonics, implemented in SCALE3 ABSPACK scaling algorithm.28 The crystal structures were solved using direct methods with SHELXT,29 and were refined by full-matrix least-squares methods on F2 with SHELXL2014. All programs used during the crystal structure determination process are included in the OLEX2 software.30 All the non-hydrogen atoms were anisotropically refined. Hydrogen atoms were introduced in calculated positions, and their coordinates were refined according to the linked atoms, with the exception of the solvent water molecules. In some cases, analysis of the F map allows to localize the water hydrogen atoms. The crystallographic details of the crystal structures are summarized in Tables S1−S3. CCDC 1827335−1827337 contain the supplementary crystallographic data for this paper. Computational Studies. The structures of HL+ and H2L2+ (L = L1; L2) were initially subjected to a conformational analysis following the procedure described by Macgregor et al.,31 and the resulting lowest energy structures were subsequently selected for optimization at the DFT level. These DFT optimizations were performed with the Gaussian 09 software package32 using the B3LYP hybrid functional33,34 and the Pople-style basis set 6-31G(d,p).35 Calculations were performed without any symmetry constraints, and the effects of the solvent were included (water, ε = 78.39) self-consistently through the polarizable continuum model (PCM).36,37 For all the structures, subsequent analytical Hessian calculations were carried out both to confirm the nature of the stationary points as minima and to derive the thermochemical corrections required to obtain free energies.

Scheme 1. Ligand Drawings

ability toward metal ions and catalytic efficiency,11−14 we have prepared a symmetrically N-alkylated derivative with two isopropyl groups and one methyl group (L1). We have conducted studies on its protonation and Cu2+ coordination behavior, both from speciation and kinetic points of view, and compared the results with those obtained for L2 and L3. The kinetic analysis is particularly interesting since we have shown that often N-alkylation may yield slower formation and dissociation kinetics. Methylation can, on the other hand, lead to a sequential dissociation mechanism with formation of observable reaction intermediates.15 Furthermore, we have made a preliminary analysis of antioxidant capability by means of the McCord−Fridovich method using nitroblue tetraazolium (NBT) as superoxide radical scavenger.16−18 The results are discussed taking into consideration the redox potentials, electron effects, hydration energies, and bulkiness of the alkylating groups.

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EXPERIMENTAL SECTION

The synthesis of 6-(N-methyl)-3,6,9-triaza-1(2,6)-pyridinecyclodecaphane,19 3,6,9-(trimethyl)-3,6,9-triaza-1(2,6)-pyridinecyclodecaphane (L2), and 3,6,9-triaza-1(2,6)-pyridinecyclodecaphane (L3) was carried out following a procedure described previously in the literature.9 All reagents were obtained from commercial sources and used as received. Solvents used for the chemical synthesis were of analytical grade and used without further purification. Synthesis of 6-(N-Methyl)-3,9-(N-isopropyl)-3,6,9-triaza-1(2,6)-pyridinecyclodecaphane Hydrochloride (L1). 6-(N-Methyl)-3,6,9-triaza-1(2,6)-pyridinecyclodecaphane trichlorohydrate (0.400 g, 0.86 mmol), 2-bromoisopropane (0.646 mL, 6.9 mmol), Na2CO3 (0.638 g, 6.02 mmol), and tetrabutylammonium bromide (0.002 g, 0.006 mmol) were suspended in dry CH3CN (10 mL), and the mixture was left to reflux (∼82 °C) under a N2 atmosphere for 48 h with stirring. After cooling, the Na2CO3 was removed by vacuum filtration and the solvent was evaporated to dryness. The crude product was purified by silica column chromatography using a mixture of CH2Cl2/methanol to afford the product as a light brown oil. The product was taken into dry ethanol (2 mL) and acidified with HCl/ dioxane until pH 1. The suspension was then centrifuged to obtain the pure product as the corresponding hydrochloride salt (0.198 g, 42%). Characterization data of the compound are included in the Supporting Information (Figures S1−S4). 1H NMR (300 MHz, CDCl3), δ (ppm): 8.01 (t, J = 7.8 Hz, 1H), 7.53 (d, J = 7.8 Hz, 2H), 4.83 (s, 2H), 4.59 (s, 2H), 3.96 (hept, J = 6.7 Hz, 2H), 3.55 (s, 4H), 2.71 (s, 2H), 2.43 (s, 3H), 3.05−2.00 (m, 2H), 1.48 (s, 6H), 1.46 (s, 6H). 13C NMR (75.43 MHz, CDCl3) δ (ppm): 150.45 140.57, B


Inorganic Chemistry

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To obtain improved values, all energies were recomputed by singlepoint calculations with the larger basis set system 6-311+G(2d,2p). Thus, energies shown herein refer to Gibbs free energies in solution, obtained by adding zero-point and thermal effects at 298.15 K and 1 atm, as well as D3(BJ) dispersion effects,38,39 to the B3LYP/6311+G(2d,2p)(PCM)//B3LYP/6-31G(d,p)(PCM) electronic energies. The Molden40 and Pymol41 packages were used for the modeling analysis. UV−vis Measurements. Water was twice distilled and passed through a Millipore apparatus. The pH values were measured with a Metrohm 713 pH meter, and adjustments of the hydrogen ion concentration of the solutions were made with diluted HCl and NaOH solutions. UV−vis absorption spectra were recorded on an Agilent 8453 spectroscopy system. Electrochemical Measurements. Cyclic voltammetry experiments were performed with 10−3 M aqueous solutions of L1, L2, and L3 at pH = 7.4. For the study of the electrochemistry of the metal complexes, equimolar amounts of Cu(ClO4)2·6H2O and the ligand were dissolved in 50 mM TRIS, previously degassed with nitrogen for 10 min, and then the voltammograms were recorded. Electrochemical experiments were performed with a BAS CV 50W and Metrohm PGSTAT 101 Autolab in a conventional threecompartment cell with a glassy-carbon working electrode. Prior to the series of experiments, the working electrode was cleaned and activated. Before each run, the electrode was polished with an aqueous suspension of alumina on a soft surface, dried, and cleaned. AgCl (3 M NaCl)/Ag and a platinum-wire auxiliary electrode completed the three-electrode configuration. The cyclic voltammograms were recorded at scan rates of 10−100 mV/s. The pH was adjusted to 7.4 by adding appropriate amounts of aqueous HCl and/ or NaOH 0.1 M solutions. Kinetic Experiments. Depending on the time scale of the reaction, the kinetic measurements were carried out either with an Applied Photophysics SX17MV stopped-flow instrument provided with a PDA-1 diode array detector or with a Cary 50-Bio spectrophotometer equipped with a thermostated multicell holder accessory. In both cases, the spectral changes over a selected wavelength range were measured and fitted to different kinetic models with the SPECFIT program,42 which provides the values of the rate constants and the electronic spectra of the different species for the model proposed. The kinetic studies of complex decomposition were carried out under pseudo-first-order conditions of acid excess, and the solutions contained Cu2+ and the corresponding ligand in a 1:1 molar ratio. The pH was adjusted using NaOH and HClO4 solutions. In all cases, the equilibrium speciation curves were used to select the starting pH of the solution so that it contains only a major complex species. Experiments with the L1 complex were carried out at 298.1 ± 0.1 K using 0.15 M NaClO4 as supporting electrolyte. However, because of the higher stability of the L2 complex, the concentrations of acid required to achieve complete decomposition of the [CuL2]2+ complex are larger than for the case of L1, and the kinetics experiments were carried out using 1 M NaClO4 as supporting electrolyte. For kinetic studies on complex formation, a 5.0 × 10−4 M solution of the ligand whose pH had been previously adjusted with HClO4 and/or NaOH solutions was mixed in the stopped-flow instrument with a solution at the same pH containing the same concentration of Cu2+. Pseudo-first-order conditions were avoided to make complexation slow enough to be measured and also to avoid complications caused by the possibility of formation of complexes with different stoichiometries. Experiments were carried out at 298.1 ± 0.1 K using 0.15 M NaClO4 as supporting electrolyte at several values of the initial pH for the metal and ligand solutions (pH range of ca. 2.0−5.5). Under these conditions, complex formation can be easily monitored, and Cu2+ initially exists almost exclusively as the aqua-complex, thus facilitating the interpretation of the kinetic data because there is not any contribution of Cu2+ hydroxo complexes to the rate of complex formation.

Article

RESULTS AND DISCUSSION

Acid−Base Behavior. The knowledge of the acid−base behavior of polyamine type ligands is crucial for understanding their coordination behavior toward metal ions. For this reason, we have used a variety of experimental and computational techniques to determine not only the values of the protonation constants of the L1−L3 ligands but also to get information about its protonation sequence. Table 1 shows the protonation constants of L1, L2, and L3 determined by pH-metric titrations in 0.15 M NaClO4 at 298.1 K.

Table 1. Stepwise Protonation Constants for L1, L2, and L343 Measured at 298.1 K in 0.15 M NaClO4 reactiona H + L ⇄ HL H + HL ⇄ H2L H + H2L ⇄ H3L a

L1 b

11.25(1) 7.345(7)

L2

L3

10.88(1) 7.37(1)

10.54(1) 7.96(1) 1.90(1)

Charges omitted. bValues in parentheses are standard deviation.

L1 and L2 display only two measurable stepwise protonation constants, while the nonalkylated L3 ligand43 has a third detectable constant in the 2.0−11.0 pH range of study. The first protonation constant is in all cases very high, its value increasing with the bulkiness of the substituents at the aliphatic nitrogens. The values of the second stepwise protonation constant are similar for L1 and L2, increasing however significantly for L3. It has been previously reported that Nmethylation decreases the basicity of both open-chain and macrocyclic polyamines (see for, example, ref 12). In the present case, the basicity constant of L1 increases with respect to that of nonmethylated L3 in the first protonation step but decreases in the subsequent two steps. At this point, it is interesting to note that, whereas open-chain polyamines as tren and trien show a decrease of the basicity for all the protonation steps upon methylation, in the case of cyclam, there is a decrease for the first two steps and an increase for the last two protonations.13,44,45 For a related scorpiand ligand, we have recently shown that the first protonation bears a small increase in basicity upon methylation.14 All of these observations reflect the importance of internal hydrogen bonds within macrocyclic structures. In the case of L1 and L2, the increase in basicity observed for the first protonation upon alkylation suggests that the proton in the HL+ species is stabilized by internal hydrogen bonding, which is supported by X-ray crystallography and by DFT calculations (see below). The acid−base behavior has been analyzed using UV and NMR spectroscopies. The variation with pH of the absorbance of the pyridine band at 226 nm shows a significant decrease in intensity from pH 9 to 7 as the diprotonated H2L12+ species is formed (Figure 1). On the other hand, the 1H NMR signal of the central methyl group shifts markedly upfield from pH 12.5 to pH 10.2 (Figure S5). However, this signal experiences an opposite downfield shift when the pH is decreased from 9 to 4. These shifts may be interpreted assuming that the first protonation occurs at the central nitrogen atom of the chain,46 and the high value of the protonation constant may be ascribed to hydrogen bonding between the protonated central amine group and the nitrogen atom of the pyridine ring. Addition of the second proton would produce a reorganization of the protons to become both placed at the nitrogen atoms functionalized with the isopropyl groups. Moreover, below C


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Inorganic Chemistry

Figure 1. Distribution diagram of L1 as a function of the pH in aqueous solution. The UV−vis spectroscopic parameters at 226 nm (red dots) are overlaid.

Figure 2. Structure of the cation [HL1]+ determined by single-crystal X-ray diffraction.

pH 9, the methyl groups of the isopropyl fragments show an upfield shift in the 1H NMR spectra supporting the protonation of the nitrogen atoms attached to them at this pH. The analysis of the variation with pH of the 1H and 13C NMR spectra of L2 gives rise to similar conclusions with the first protonation occurring at the central amine group of the polyamine, and the second one locating both protons at the amine groups closer to the pyridine ring (see Figure S6). The observations in aqueous solution are in agreement with the crystal structure of the monoprotonated salt of L1. The asymmetric unit of [HL1]2(ClO4)2·H2O (space group P21/n) consists of two very similar [HL1]+ cations, perchlorate counteranions, and one water molecule. The three alkyl groups are located in cis with respect to each other, and the proton is bound to the central amine group as inferred from the solution studies. The protonated amine is hydrogen bonded in both cations to the pyridine nitrogen, being the N−H···N distances 1.898 and 1.845 Å and the angles 151.45° and 152.29°, respectively (Figure 2). This hydrogen bond should somehow force the cis-disposition adopted by the alkyl substituents. The angle of the pyridine ring with the plane passing through the macrocyclic amines is ca. 20°. DFT calculations were also conducted to find out the most stable conformers formed at each protonation step and to estimate the energetic barriers separating them. The calculations show, in agreement with the crystal structure and solution studies, that, in the most stable monoprotonated species, the proton is placed in the central amine (N3) of the macrocyclic cavity, while, in the diprotonated species, both protons are bound to the lateral amines substituted with the isopropyl groups (N2 and N4; see Figure S7). Interestingly,

the most stable conformers of the mono- and diprotonated species show the alkyl groups oriented in cis with respect to the N1/N2/N4 plane, as observed in the crystal structure of the monoprotonated species. The additional optimized conformers of HL1+ and H2L12+ are included in Figures S8 and S9. The analysis of the HL1+ conformers shows that those structures where the rotation of the isopropyl groups at N2 and N4 has taken place are up to ca. 3 kcal/mol higher in energy, whereas conformers where the inversion of N2 (or N4) has taken place are at least 4−5 kcal/ mol less stable. In addition, conformers where the proton is not located at N3 but at N2 or N4 start to appear at 5−7 kcal/ mol relative free energies. In the case of H2L12+, it is again observed that structures featuring different orientations of the isopropyl groups are all within a 4 kcal/mol free energy range. However, the presence of the additional proton seems to prevent the formation of low energy conformers where the inversion of the N2 center has taken place, as the first conformer with this feature appears at a relative free energy of 8.2 kcal mol−1. Additionally, alternative proton arrangements are also possible for H2L12+, these being at least ca. 7 kcal/mol less stable than the form with the protons located at nitrogens N2 and N4. The absence of isopropyl groups in HL2+ and H2L22+ simplifies the conformational analyses. It is nonetheless interesting to observe that the substitution of isopropyl by methyl groups in L2 results in a stabilization of the structures where the inversion of the N2 center has taken place, which appear 0.9 and 4.2 kcal mol−1 above the most stable structures of HL2+ and H2L22+, respectively (Figures S10 and S11). This indicates that L2 is more flexible than L1, which is expected to D


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Inorganic Chemistry

to reduction in the number of M-N-H···O hydrogen bonds with the water solvent molecules and increase in the radii of the complexes.13,46 Although the values of the stability constants indicate a lower stability of the complexes with the alkylated ligands, to make straight through comparisons, the different affinities of the ligands for protons have to be balanced. To do so, the overall percentages of complex formation for solution mixtures containing 1 equiv of copper and 1 equiv of each ligand have to be compared. The graphic in Figure S12 shows that, under those conditions, 80% of the Cu2+ will be complexed by L3 and the rest by L2, while the amount of Cu2+ complexed by L1 would be negligible. Nevertheless, this lower stability of [CuL1]2+ may help the cycling between oxidation states II and I of copper required for the dismutation of superoxide to occur. Two different crystalline samples of the system Cu2+-L1 were obtained by slow evaporation of aqueous solutions containing Cu(ClO4)2·6H2O and L1·3HCl with and without an excess of NaCl. Crystals of formula [Cu(L1)Cl](ClO4) (1) evolved in the first case, while, in the second case, the chloride ligand was replaced by a water molecule, giving rise to crystals of formula [Cu(L1)(H2O)](ClO4)2·H2O (2) (see Figure 4). The coordination geometry around the copper center is square pyramidal in both structures, with the central nitrogen atom of the polyamine chain occupying the slightly elongated axial position. The Cu−pyridine nitrogen is the shortest distance in both crystal structures (see selected bond distances and angles in Tables S2 and S3). The macrocycle adopts a more bent conformation in the complexes than in the noncoordinated ligand; the angle defined by the planes of the pyridine ring and

facilitate the geometrical rearrangements required to accommodate a metal center with a lower energy cost. Interaction with Cu2+ Ions. The interaction of Cu2+ with L1 and L2 has been studied by pH-metric titration, UV−vis and, in the case of L1, by X-ray diffraction studies. The stability constants for L1, L2 and those reported for the reference ligand L35,43 are collected in Table 2, whereas the distribution diagrams are shown in Figure 3. Table 2. Logarithm of the Stability Constants for the Cu2+ Complexes of L1, L2, and L343 Determined in 0.15 M NaClO4 at 298.1 reactiona

L1

L2

L3

Cu + L ⇄ CuL CuL + H2O ⇄ CuL(OH) + H CuL(OH) + H2O ⇄ CuL(OH)2 + H

14.04(1)b −8.11(3) −10.04(8)

16.44(3) −8.53(4) −11.31(8)

17.78(2)

a

Charges omitted. bValues in parentheses are standard deviations in the last significant figure.

The speciation of these systems is simple, with formation of [CuL]2+, monohydroxylated and dihydroxylated mononuclear species for both L1 and L2. The stability constants of the [CuL]2+ complexes follow the trend L3 > L2 > L1, thus showing that the bulkier and more hydrophobic the substituents, the smaller the stability constants are. A similar decrease of the stability of the complexes associated with alkylation of secondary amine groups has been previously reported and explained in terms of a poorer σ-donor ability due

Figure 3. Species distribution curves for the Cu2+-L1 (a) and the Cu2+-L2 complexes (b) in 0.15 M NaClO4. E


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Inorganic Chemistry

Figure 4. View of the crystal structure of the Cu2+ complexes: (a) [Cu(L1)Cl](ClO4), (b) [Cu(L1)(H2O)](ClO4)2·H2O.

As found for related polyamine complexes, addition of an acid excess to solutions containing the Cu2+ complexes of L1 and L2 results in complex decomposition according to eq 1, and the kinetics of the process can be easily monitored by measuring the disappearance of the absorption band at ca. 700 nm. The decomposition of the [CuL1]2+ and [CuL2]2+ complexes was found to be a relatively slow process that can be monitored with conventional UV−vis spectrophotometry. The fit of the spectral changes to a single exponential was satisfactory in all cases and yielded kobs values that show a linear dependence with respect to the acid concentration (eq 2; Figures 6 and S14), with second-order rate constants k =

that one going through the polyamine nitrogen atoms opens up to reach values of 73.3° for 1 and 82.2° for 2. The ancillary ligands, Cl− in 1 and H2O in 2, occupy the equatorial position in trans with respect to the pyridine nitrogen, thus being placed in a hydrophobic domain surrounded by the alkyl groups. While in 1 the complex units are isolated, in 2, each complex unit connects with another one through a hydrogen bond network involving the coordinated water molecule, two bulk water molecules, and the perchlorate counteranions (see Figure S13). Previously to perform the kinetic studies on the interaction between Cu2+ and the ligands, the spectra of solutions containing Cu2+ and the corresponding ligand in a 1:1 molar ratio were recorded after reaching equilibrium at different pH values selected from the species distribution curves in Figure 3. These spectra are included in Figure 5, and they show that the

Figure 6. Plot of the dependence on the acid concentration of the observed rate constants for the acid-promoted decomposition of solutions containing [CuL1] 2+ (black circles) and [CuL1(OH)]+(white circles) complexes and a mixture of both species (triangles) in 0.15 M NaClO4. For [CuL1(OH)]+, there is a previous kinetic step too fast to be measured even with the stopped-flow technique.

Figure 5. Electronic spectra for aqueous solutions containing the [CuL1(OH)]3+ (orange dashed line), [CuL1]2+ (green dashed-dotted line), and [CuL2]2+ (red solid line) complexes at 298.1 ± 0.2 K in the presence of 0.15 M NaClO4 (C0Cu = C0L= 1 × 10−3 M).

(1.96 ± 0.02) × 10−4 M−1 s−1 for [CuL1]2+ and k = (4.61 ± 0.17) × 10−3 M−1 s−1 for [CuL2]2+. These results indicate that the introduction of the bulkier isopropyl groups causes a deceleration of complex decomposition by a factor of ca. 20. On the other hand, kinetic experiments on the decomposition of the [CuL1(OH)]+ species revealed that it is converted to [CuL1]2+ within the stopped-flow mixing time (ca. 1.7 ms for the instrument used) and then decomposition is completed with similar kinetics to [CuL1]2+. The rate law in eq 2 can be interpreted in terms of the general mechanism proposed initially by Margerum, and refined by others, for the acidpromoted dissociation of polyamine complexes. According to that mechanism, one nitrogen dissociates from the metal ion without its replacement by a water molecule, resulting in the formation of an activated intermediate susceptible to bear parallel attacks by the solvent or by a proton.47−49 In the

formation of Cu2+-L complexes is signaled by the appearance of an absorption band at 705 nm for [CuL1]2+ and 690 nm for [CuL2]2+. These values are close to that found for the related [CuL3]2+ complex (695 nm),43 in agreement with a similar coordination environment of the metal ion in the three complexes. In contrast, the [CuL1(OH)]+ complex shows a shift in the band (750 nm) that can be ascribed to the stronger π-donor character of the hydroxo ligand. Kinetics of Formation and Decomposition of Cu2+ Complexes. To complement the information on the Cu2+ complexes, kinetic studies on complex formation and dissociation were also carried out. For simplicity, the results obtained for the acid-promoted dissociation of the metal ion will be presented first. F


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Inorganic Chemistry

was made, and the values of the observed rate constants were found to change with pH, in agreement with the existence of fast acid−base equilibria previous to rate-limiting complex formation. However, as no buffering agent was used, the proton concentration changes during the course of the reaction and this could lead to changes of kobs with pH. For that reason, the data were then fitted to the model proposed in eqs 4−6, which includes all relevant acid−base pre-equilibria for ligand protonation, followed by rate-determining complex formation through reaction of Cu2+ with H2L2+, the major ligand species under the experimental conditions used. The equilibrium constants for ligand protonation were fixed at the values derived from potentiometric studies, and the proton concentration was allowed to change with time, but the kobs values so derived were still found to change with the starting pH (see Figures 7 and S15). The plots in these figures suggest

present case, the absence of a significant nonzero intercept indicates that complex decomposition occurs exclusively through the pathway involving proton attack. At this point, it is interesting to compare the present kinetic data with those previously reported for the complex with the related unalkylated macrocycle L3.43 Although [CuL3]2+ decomposes with a second-order dependence with respect to the acid concentration, which indicates that the ratedetermining step is shifted to the breaking of the second Cu−N bond, the values of kobs for [H+] values similar to those used in the present case for the L1 and L2 complexes are about 2 orders of magnitude faster for the L1 complex, which is surprising given the similar structure and spectra of the three complexes. However, the present values for [CuL1]2+ and [CuL2]2+ are closer to those found for decomposition of [CuHL]3+ species with related macrocycles containing a protonated pendant arm, which also decompose with firstorder kinetics with respect to the proton concentration.50 In that case, the kinetic changes with respect to the unsubstituted macrocycle L3 were interpreted in terms of the classical mechanism commented above and were associated with the different structural distortions induced by the presence of the protonated donor group in the pendant arm. The observation of similar kinetic effects for [CuL1]2+ and [CuL2]2+ suggests that the deceleration of the decomposition process is probably more related to the effect of alkylation than to protonation of the uncoordinated donor group. [CuL]2 + + H+exc → Cu 2 + + HxLx +

(1)

kobs = k[H+]

(2)

Figure 7. Plot of the dependence on the pH of the observed rate constant for the formation of the [CuL1]2+ complex at 298.1 K in the presence of 0.15 M NaClO4.

2+

The kinetics of complex formation of Cu with L1−L2 was studied in moderately acidic media under non-pseudo-firstorder conditions by mixing solutions containing the metal ion and the ligand in a 1:1 molar ratio. The experiments were carried out within a pH range of ca. 2.0−5.5 without adding any buffering agent because preliminary experiments revealed a significant effect of the buffer on the kinetics. Actually, it has been previously shown that the addition of buffers often introduces some complications in this kind of kinetic studies probably due to the interaction between the buffer and the metal ion or the protonated ligand species.51−55 Despite the close similarity between L1 and L2, the kinetics of complexation with both ligands are significantly different. Whereas, for L1, the kinetics could be monitored with a conventional UV− vis spectrophotometer, a stopped-flow instrument had to be used for L2 because of the faster kinetics. Thus, as observed for complex decomposition, it appears that the steric effects associated with the presence of the two i-Pr groups also cause a decrease in the rate of complex formation. For both ligands, complex formation (eq 3) occurs in a single kinetic step and the final spectra agree well with those of the corresponding [CuL]2+ complexes. It is important to note that complex formation is observed for both ligands at pH values as low as 2. Whereas, for L2, this finding is in agreement with the equilibrium results, which anticipate formation of significant amounts of [CuL2]2+ at those pH values, the formation of [CuL1]2+ under those conditions suggests that the equilibrium constant for complex formation in Table 2 is somewhat underestimated, surely because of the extremely slow kinetics of complexation under strongly acidic conditions. A preliminary fit of the data to a single second-order process, first-order with respect to both the metal ion and the ligand,

the occurrence of an additional protonation pre-equilibrium with a log K value close to 2.5−3.5. However, inclusion of a hypothetical H3L3+ species in the kinetic model leads to unsatisfactory results because it leads to log KH3L values of 2.6 (L1) and ca. 3 (L2) with erratic changes in the values of kCuL. Moreover, the values derived for log KH3L are large enough for making those protonation processes observable experimentally, and they are not observed in the potentiometric titrations. Cu 2 + + HxLx + → [CuL]2 + + x H+

(3)

L + H+ ⇆ HL+;

(4)

KHL

HL+ + H+ ⇆ H 2L2 +;

K H 2L

(5)

Cu 2 + + H 2L2 + → [CuL]2 + + 2H+;

k CuL

(6)

To avoid any possibility that the observed pH dependence of the rate constants was caused by the changes in the proton concentration during the kinetic runs, we decided to replace the kobs values in Figures 7 and S15 by those estimated using the method of the initial rates. In this way, the plots in Figures S16 and S17 showing the changes of the rate constant with pH were obtained. Those plots again showed the occurrence of an acid−base equilibrium that cannot be assigned to ligand protonation and so we considered the possibility that the pH dependence of the rate of complex formation is caused by an acid−base equilibrium for a reaction intermediate, so that complexation occurs through a mechanism of the type shown in Figure 8. In this model, a pre-equilibrium involving the HL+ G


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Inorganic Chemistry a=

k 3k5KH k −3 + k 5

(8)

b = K1K 2k4 c= k1 =

and H2L2+ species is considered, and the fully deprotonated species (L) is ignored for simplicity because its concentration is negligible under all experimental conditions used. The model considers that Cu2+ reacts with both HL+ and H2L2+ to form the reaction product through two parallel pathways. Reaction with H2L2+ leads to the sequential formation of intermediates [CuH2L]4+ and [CuHL]3+A, which interconvert through an acid−base equilibrium. Formation of the final [CuL]2+ complex is considered to occur through reorganization of intermediate [CuHL]3+A. The pathway involving HL+ implies formation of an intermediate [HCuL]3+B, which is considered to have a structure different from [CuHL]3+A because the nitrogen atoms which are protonated are different in the HL+ and H2L2+ species. As the reaction occurs in a single kinetic step without accumulation of any reaction intermediate, all intermediates can be considered to be formed under steadystate conditions, which leads to the rate law in eq 7 with the values of a, b, and c in eqs 8−10. Equation 10 includes a simplification of the expression for c which is justified because it must be reasonably expected that k2 ≫ k−1 because k2 corresponds to proton dissociation from [CuH2L]4+ and k−1 corresponds to the presumably slower Cu2+ dissociation from the same intermediate. Although all the three parameters include contributions from different rate and equilibrium constants, the equations clearly indicate that a measures the contribution of the HL+ pathway, whereas b and c correspond to the contribution of the H2L2+ pathway. In addition, the value of the mechanistically relevant rate constant k1 can be obtained by using eq 11. (a + b)[H+] + ac ([H+] + KH)([H+] + c)

k4(k −1 + k 2) kk ≈ 4 2 k −1k −2 k −1k −2

(10)

b c

(11)

An adequate fit of the data in Figures S16 and S17 to eq 7 was obtained, and the values of a, b, c, and k1 are shown in Table 3. The faster formation of the L2 complex is the result of a combined effect of larger values of a and b, but the most important finding is the negligible value of a for L1, which means that, for this ligand, metal complexation occurs exclusively through reaction of the metal ion with H2L12+. In contrast, for L2, both pathways involving reaction with HL2+ and H2L22+ are operative. As the concentrations of the HL+ species under the experimental conditions used in the kinetic experiments are not very different, the absence of an operative pathway involving HL1+ indicates that the steric effects associated with the presence of the isopropyl groups prevents the formation of species able to evolve into the observed CuL complexes. With regard to the nature of intermediates initially formed in both pathways, [CuH2L]4+ and [CuHL]3+B, one reasonable possibility is that they correspond to species with the ligand acting as monodentate. These species are expected to be formed via the interaction of the HOMO orbital of the mono- or diprotonated ligand and the LUMO of an aqueous Cu2+ species, and therefore, the composition of these ligand orbitals can provide information about these intermediates. Analysis of the HOMO orbitals of the diprotonated H2L12+ and H2L22+ ligands (see Figure 9) shows that these are mostly located at their tertiary N3 central amine, with the presence of methyl or isopropyl ligands at the N2 and N4 atoms not interfering in the formation of Cu−N3 bonds. On the contrary, the HOMOs of HL1+ and HL2+ (see Figure 9) show that these are located at the lone pairs of N2 and N4, pointing toward the center of the macrocycle cavity. Thus, when the approach of Cu2+ is considered, the presence of i-Pr groups in HL1+ is expected to cause a larger steric effect than the Me groups in HL2+. Although the whole mechanistic picture is, evidently, much more complicated than this, the orbital analysis of L1 and L2 is in agreement with the negligible participation of the HL1+ pathway and the more important role of this pathway in the case of L2. On the other hand, the values of k1 in Table 3 also deserve some additional comment. The values for L1 and L2 are significantly larger than that previously reported for L3. In the cases of L1 and L2, although the rate-determining step for the formation of the final complex corresponds to a later stage, k1 corresponds to the rate constant for the formation of Cu2+-

Figure 8. Proposed mechanism for the formation of [CuL]2+ complexes.

kobs =

(9)

(7)

Table 3. Summary of Kinetic Parameters Derived from the Fit of the Kinetic Data in Figures S16 and S17 to eq 12 a (s−1) b (s−1) c (M) k1 (M−1 s−1)

L1

L2

L3

(1.8 ± 0.2)×10−4 (8 ± 1)×10−4 0.22 ± 0.04

(4.05 ± 0.08)×10−4 (7.9 ± 0.4)×10−3 (1.7 ± 0.1)×10−4 46 ± 4

0.011 ± 0.002a

a

Value from ref 1. H


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Inorganic Chemistry

present k1 values with respect to the rate constant for water exchange in Cu2+aq cannot be justified exclusively on the basis of the electrostatic effects introduced by the repulsion between the positively charged metal ion and H2L2+ species because the Kos value for the formation of an outer-sphere complex between two species with charge +2 is of the order of 7.4 × 10−3,58 which should lead to k1 values several orders of magnitude faster than those observed experimentally. Thus, it is evident that the present small k1 values reflect the steric requirements of the H2L2+ species, in agreement with previous observations by Rorabacher, that established that the rate of complexation decreases with alkylation in the order NH3 > RNH2 > R2NH > R3N and that changing the nature of the R group from methyl to ethyl causes an additional deceleration.59 If the first Cu−N bond is considered to occur with the central NMe group in the aliphatic chain, the different values of k1 observed for L1 and L2 indicate an additional effect of the substituents at the adjacent nitrogens. In the case of L3, a resolved value of k1 is not available. However, the lower steric requirements of L3 and the higher stability of the [CuL3]4+ complex anticipate a value of k1 higher than those observed for L1 and L2, and so the experimental observation of a rate constant of 0.011 M−1 s−1 for complexation between Cu2+ and H2L32+ would clearly indicate that the rate-determining step does not correspond to formation of the first Cu−N bond. Electrochemistry. In deoxygenated solutions, the voltammograms of the Cu(II)-L3 (Figure 10) solutions display a weak cathodic peak at ca. 0.0 V (C1), followed by a much more intense signal at −0.50 V which looks like two consecutive, highly overlapped peaks (C21, C22). In the subsequent anodic scan, a typical stripping peak was recorded at −0.05 V (ACu), followed by a shoulder at 0.10 V (A1). Upon shifting the switching potential to values near to −0.50 V, the stripping peak vanishes so that, finally, only the extremely weak C1/A1 signals remain, appearing as an essentially reversible one-

Figure 9. HOMO of the most stable mono- and diprotonated forms of L1 and L2 (isovalue = 0.07).

H2L2+ species with the ligand acting as monodentate, so that they represent a direct measure of the rate of formation of the first Cu−N bond. As expected, its value decreases when the steric requirements of the ligand increase, as observed also for the overall stability constant of the complexes. This process is expected to occur through the classical Eigen-Wilkins mechanism, with initial formation of an outer-sphere complex with equilibrium constant Kos, followed by rate-determining substitution of coordinated water by the entering nitrogen, so that the rate constant for complexation has a value close to the product Kos × kexc, where kexc is the rate constant for water exchange in the metal aqua ion,56 which is of the order of 4.4 × 109 s−1 for Cu2+aq.57 The extremely large decrease of the

Figure 10. Cyclic voltammograms at the glassy carbon electrode of a 1.0 mM solution of Cu2+-L3 and Cu2+-L1 (inset) in 0.15 M NaClO4 at pH 7.0. Potential scan initiated at 0.25 V in the negative direction and switched at −0.25 (red line), −0.45 (ochre line), and −0.65 V (blue line) vs Ag/ AgCl. Potential scan rate 50 mV s−1. I


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Inorganic Chemistry

Figure 11. Cyclic voltammograms at glassy carbon electrode of air-saturated (red line) and deoxygenated (blue line) 1.0 mM CuIL1 solutions in 0.15 M NaClO4 (pH 7.0). Potential scan initiated at 0.25 V in the negative direction; potential scan rate 50 mV s−1.

2Cu IL → Cu IIL + Cu + L

electron couple, as judged by the cathodic-to-anodic peak potential separation tending to 60 mV at low scan rates (see the Supporting Information, Figure S18). Accordingly, the overall reductive process consists of the reduction of CuIIL to Cu metal, the intermediate CuIL species being rapidly reduced. A similar behavior was observed for L2, whereas, for L1, the CuIL species was sufficiently stabilized under electrochemical conditions to display a well-defined couple C21/A21 at a midpeak potential of −0.30 V, as can be seen in the inset in Figure 10. This voltammetry can be interpreted in terms of a square scheme (see the Supporting Information, Figure S19) involving the reduction of the CuIIL complex to an analogue CuIL form (C2/A2 couple) coupled to dissociation resulting in the reduction of a mixture of CuII-chloride-complexes to CuICly ones (C1/A1 couple), also accompanied by the rapid and irreversible reduction of CuIL to Cu(0) through the process C22. These results are consistent with the recognized stabilization of the Cu(I) oxidation state upon alkylation of macrocyclic polyamines, studied by Meyerstein et al.13 The formal potential of the CuIIL/CuIL couple for L1 (−0.30 V vs Ag/AgCl) is consistent with the appearance of SOD activity, but the similarity in the voltammetric profiles of all three complexes suggests that the differences between them derive mainly from kinetic factors. Such differences can be attributed to the possibility of a disproportionation reaction of the electrochemically generated CuIL complex. Thus, in all cases, the parent CuIIL species experiences an essentially reversible one-electron reduction Cu IIL + e− → Cu IL

For L2 and L3, the disproportionation is fast and the voltammetric profile corresponds, essentially, to that of a twoelectron reduction process (see Figures 10 and 11). However, in the case of L1, possibly due to its larger steric constraints, the disproportionation process (eq 14) is relatively slow and the intermediate CuIL species is stable enough to display a well-defined CuIIL/CuIL couple (see Figure 11). In the presence of dissolved O2, the voltammetry changes significantly. In the absence of copper complexes, O2 is reduced (oxygen reduction reaction, ORR) through an apparently irreversible process at ca. −0.65 V (Cox). This reduction has been described60 as a complicated multistep reduction initiated by the one-electron reduction yielding superoxide radical anion O2 + e− → O2•− (ads)

(15)

followed by hydration competing with superoxide disproportionation, yielding hydroperoxide (HO2·) and hydroxyl (HO·) radicals, resulting in the formation of H2O2.60 In the presence of the studied L2 and L3 copper complexes, the signals C2 and Cox merge to a unique wave retaining the stripping signal for the oxidative dissolution of metallic copper (see Figure S20 in the Supporting Information), but in these cases, the intensity of the merged cathodic wave is only slightly larger than the sum of the peak currents of Cox and C2 processes measured separately, thus suggesting that, if existing, there are no important electrocatalytic effects to be accounted. A remarkably different situation was observed for L1, for which a significant enhancement of the cathodic signals was recorded, as depicted in Figure 11. It is interesting to note that the profile of the prominent reduction signal changes from the initial cathodic scan in which it looks like two superimposed signals at −0.4 and −0.6 V (marked by * in the figure), to the second scan, where it

(12)

followed by the irreversible reduction to Cu(0) at more negative potentials Cu IL + e− → Cu + L

(14)

(13)

which competes with the disproportionation reaction. J


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Inorganic Chemistry

with pH and DFT calculations support a change in the location of the protons when going from HL+ to H2L2+ species, and these conclusions are further supported by the crystal structure of the [HL1]+ cation. With respect to the Cu2+ complexes, the introduction of isopropyl groups (L3) leads to lower stability and slower kinetics of complex formation. In addition, the analysis of the kinetics of complexation reveals a mechanism in which the acid−base properties of the intermediates initially formed determine the pH dependence of the net rate of complexation and provides a direct measure of the rate of formation of the first Cu−N bond despite the fact that the overall complexation process is kinetically controlled by formation of one of the subsequent bonds. The CVs also reflect the higher inertness of complexes with this ligand, which allows for a higher stabilization of the Cu(I) oxidation state that can be related to the higher SOD activity. Although the results must be considered only preliminary, the higher stability of the Cu(I) complexes would facilitate the Cu(II)/ Cu(I) cycling required for catalyzing superoxide dismutation.

appears that the copper signals are again obtained but followed by an enhanced oxygen reduction wave (marked by **). Among other possibilities, a possible scheme consists on assuming that the one-electron initial oxidation of CuIIL produces a CuIL species (eq 12) which acts catalytically on the first step of O2 reduction (eq 15) Cu IL + O2 → Cu IIL + O2•−

(16)

generating superoxide anion which is rapidly protonated forming the hydroperoxide radical. Under electrochemical conditions, the disproportionation of the electrochemically generated CuIL complex (eq 14) would be faster than reaction 16 in the cases of L2 and L3 and no significant electrocatalytic effects appear. In the case of L1, however, the disproportionation rate would be considerably lowered, thus providing opportunity for the appearance of catalytic effects on the ORR process. Then, the accessibility of the CuIIL/CuIL couple permits the possibility of mimicking SOD activity through HO2• + Cu IIL + OH− → Cu IL + O2 + H 2O

(17)

HO2• + Cu IL → Cu IIL + HO2−

(18)

■

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.8b01492. Complete characterization of L1, distribution diagrams, DFT-optimized structures, crystallographic data, and deduction of eq 7 (PDF)

Accordingly, the significant differences in SOD activity (vide infra) would be ultimately related, under electrochemical conditions, to the kinetics of the disproportionation reaction in turn sensitive to the different steric constraints imposed by the receptors. SOD Activity. The enzymatic assay developed by McCord−Fridovich was used to obtain the IC50 and the catalytic constants of the Cu2+ complexes at physiological pH.16−18 Table 4 shows the values obtained for the Cu2+ complexes of L1, L2, and L3, as well as some examples reported from the references.

Accession Codes

CCDC 1827335−1827337 contain the supplementary crystallographic data for this paper. These data can be obtained free of charge via www.ccdc.cam.ac.uk/data_request/cif, or by emailing [email protected], or by contacting The Cambridge Crystallographic Data Centre, 12 Union Road, Cambridge CB2 1EZ, UK; fax: +44 1223 336033.

Table 4. SOD Activity of the Cu2+ Complexes with the L1, L2, and L3 Ligands Obtained by the McCord−Fridovich Testa Cu-L1 Cu-L2 Cu-L3 Cu(ClO4)2 CuZn-SOD

6

IC50/μM

kcat/10 (M

0.26(5) 2.9(6) 2.1(4) 1.1(1) 0.010(2)

13.7 1.2 1.7 2.7 430

−1

−1

s )

ASSOCIATED CONTENT

S Supporting Information *

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AUTHOR INFORMATION

Corresponding Authors

ref

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

this work this work this work ref 64 ref 64

ORCID

Á lvaro Martínez-Camarena: 0000-0001-9910-6961 Andrés G. Algarra: 0000-0002-5062-2858 Antonio Domenech-Carbó: 0000-0002-5284-2811 Manuel G. Basallote: 0000-0002-1802-8699 Enrique García-España: 0000-0002-4601-6505

a

Some examples from the literature are also included for comparison.

The first aspect that deserves to be commented is the high catalytic constant of Cu-L1, which ranks among the best values so far reported for SOD synthetic systems.61−63 The alkylation with isopropyl groups (L1) clearly increases the activity of this family of Cu2+ complexes, when compared to that of L2 or L3. The lower stability of the [CuL1]2+ complex, associated with the poorer σ-donor ability of the tertiary amino groups, and the rigidity of the system, caused by the bulky isopropyl groups, favors the fluctuation between the oxidized Cu2+ and the reduced Cu+ forms of the metal ion.

Author Contributions

The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript. Author Contributions ⊥

These authors contributed equally.

Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS Financial support by the Spanish Ministerio de Economiá y Competitividad and FEDER funds from the EU (Projects CTQ2013-48917-C3-1-P, CTQ2015-65707-C2-2-P, CTQ2016-78499-C6-1-R, and Unidad de Excelencia MDM 2015-0538) and Generalitat Valenciana (Project PROMETEO

CONCLUSIONS The results here presented reveal an important effect of the alkyl groups on the properties of small N-alkylated tetraazacyclophanes and especially of their Cu2+ complexes. Regarding the free ligands, the changes in the NMR spectra K


Article

Inorganic Chemistry

(17) Beauchamp, C.; Fridovich, I. Superoxide Dismutase: Improved Assays and an Assay Applicable to Acrylamide Gels. Anal. Biochem. 1971, 44, 276−287. (18) Zhou, J. Y.; Prognon, P. Raw Material Enzymatic Activity Determination: A Specific Case for Validation and Comparison of Analytical Methods−the Example of Superoxide Dismutase (SOD). J. Pharm. Biomed. Anal. 2006, 40, 1143−1148. (19) Meijer, A. Manganese Chelates and Their Use as Contrast Agents in Magnetic Resonance Imaging. Patent WO2011073371A1, 2011. (20) Garcia-Espana, E.; Ballester, M. J.; Lloret, F.; Moratal, J. M.; Faus, J.; Bianchi, A. Low-spin Six-coordinate cobalt(II) Complexes. A Solution Study of tris(violurato)cobaltate(II) Ions. J. Chem. Soc., Dalton Trans. 1988, 2, 101−104. (21) Fontanelli, M.; Micheloni, M. In Proceedings of the I SpanishItalian Congress of Thermodinamics of Metal Complexes; 1990. (22) Gran, G. Determination of the Equivalence Point in Potentiometric Titrations. II. Analyst 1952, 77, 661−671. (23) Rossotti, F. J. C.; Rossotti, H. Potentiometric Titrations Using Gran Plots: A Textbook Omission. J. Chem. Educ. 1965, 42, 375. (24) Gans, P.; Sabatini, A.; Vacca, A. Investigation of Equilibria in Solution. Determination of Equilibrium Constants with the HYPERQUAD Suite of Programs. Talanta 1996, 43, 1739−1753. (25) Alderighi, L.; Gans, P.; Ienco, A.; Peters, D.; Sabatini, A.; Vacca, A. Hyperquad simulation and speciation (HySS): a utility program for the investigation of equilibria involving soluble and partially soluble species. Coord. Chem. Rev. 1999, 184, 311−318. (26) Covington, A. K.; Paabo, M.; Robinson, R. A.; Bates, R. G. Use of Glass Electrodes in Deuterium Oxide and the Relation between the Standarized pD Scale and the Operational pH in Heavy Water. Anal. Chem. 1968, 40, 700−706. (27) Hope, H. Cryocrystallography of Biological Macromolecules: A Generally Applicable Method. Acta Crystallogr., Sect. B: Struct. Sci. 1988, 44, 22−26. (28) CrysAlis PRO; Agilent Technologies Ltd.: Yarnton, Oxfordshire, England, 2014. (29) Sheldrick, G. M. SHELXT - Integrated Space-Group and Crystal-Structure Determination. Acta Crystallogr., Sect. A: Found. Adv. 2015, 71, 3−8. (30) Dolomanov, O. V.; Bourhis, L. J.; Gildea, R. J.; Howard, J. A. K.; Puschmann, H. OLEX2: A Complete Structure Solution, Refinement and Analysis Program. J. Appl. Crystallogr. 2009, 42, 339−341. (31) Häller, L. J. L.; Page, M. J.; Erhardt, S.; Macgregor, S. A.; Mahon, M. F.; Naser, M. A.; Vélez, A.; Whittlesey, M. K. Experimental and Computational Investigation of C−N Bond Activation in Ruthenium N-Heterocyclic Carbene Complexes. J. Am. Chem. Soc. 2010, 132, 18408−18416. (32) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G. A.; Nakatsuji, H.; Caricato, M.; Li, X.; Hratchian, H. P.; Izmaylov, A. F.; Bloino, J.; Zheng, G.; Sonnenberg, J. L.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.; Hasegawa, J.; Ishida, M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Vreven, T.; Montgomery, J. A., Jr.; Peralta, J. E.; Ogliaro, F.; Bearpark, M.; Heyd, J. J.; Brothers, E.; Kudin, K. N.; Staroverov, V. N.; Keith, T.; Kobayashi, R.; Normand, J.; Raghavachari, K.; Rendell, A.; Burant, J. C.; Iyengar, S. S.; Tomasi, J.; Cossi, M.; Rega, N.; Millam, J. M.; Klene, M.; Knox, J. E.; Cross, J. B.; Bakken, V.; Adamo, C.; Jaramillo, J.; Gomperts, R.; Stratmann, R. E.; Yazyev, O.; Austin, A. J.; Cammi, R.; Pomelli, C.; Ochterski, J. W.; Martin, R. L.; Morokuma, K.; Zakrzewski, V. G.; Voth, G. A.; Salvador, P.; Dannenberg, J. J.; Dapprich, S.; Daniels, A. D.; Farkas, O.; Foresman, J. B.; Ortiz, J. V.; Cioslowski, J.; Fox, D. J. Gaussian 09; Gaussian, Inc.: Wallingford, CT, 2009. (33) Becke, A. D. Density-Functional thermochemistry.III. The Role of Exact Exchange. J. Chem. Phys. 1993, 98, 5648. (34) Lee, C.; Yang, W.; Parr, R. G. Development of the ColleSalvetti Correlation-Energy Formula into a Functional of the Electron

II2015-002) is gratefully acknowledged. A.M.-C. wants to thank the Spanish Ministry of Education, Culture and Sport for the Ph.D. grant FPU14/05098.

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REFERENCES

(1) Bianchi, A.; García-España, E. Azacycloalkanes and Azacyclophanes. In Supramolecular Chemistry: From Molecules to Nanomaterials; Steed, J. W., Gale, P. A., Eds.; Wiley: Hoboken, NJ, 2012; pp 753−784. (2) Fenton, D. E. The Chemistry of Macrocyclic Ligand Complexes; Lindoy, L. F., Ed.; Cambridge University Press: Cambridge, U.K., 1989; Vol. 47. (3) Bradshaw, J. S.; Krakowiak, K. E.; Izatt, R. M. The Chemistry of Heterocyclic Compounds: Aza-Crown Macrocycles; John Wiley & Sons, Inc.: Hoboken, NJ, 1993; Vol. 51. (4) Gloe, K. Macrocyclic Chemistry; Springer-Verlag: Berlin, 2005. (5) Costa, J.; Delgado, R. Metal Complexes of Macrocyclic Ligands Containing Pyridine. Inorg. Chem. 1993, 32, 5257−5265. (6) Lincoln, K. M.; Richardson, T. E.; Rutter, L.; Gonzalez, P.; Simpkins, J. W.; Green, K. N. An N-Heterocyclic Amine Chelate Capable of Antioxidant Capacity and Amyloid Disaggregation. ACS Chem. Neurosci. 2012, 3, 919−927. (7) Lincoln, K. M.; Gonzalez, P.; Richardson, T. E.; Julovich, D. a.; Saunders, R.; Simpkins, J. W.; Green, K. N. A Potent Antioxidant Small Molecule Aimed at Targeting Metal-Based Oxidative Stress in Neurodegenerative Disorders. Chem. Commun. 2013, 49, 2712−2714. (8) Serrano-Plana, J.; Aguinaco, A.; Belda, R.; Garcia-España, E.; Basallote, M. G.; Company, A.; Costas, M. Exceedingly Fast Oxygen Atom Transfer to Olefins via a Catalytically Competent Nonheme Iron Species. Angew. Chem., Int. Ed. 2016, 55, 6310−6314. (9) Serrano-Plana, J.; Oloo, W. N.; Acosta-Rueda, L.; Meier, K. K.; Verdejo, B.; García-España, E.; Basallote, M. G.; Münck, E.; Que, L.; Company, A.; Costas, M. Trapping a Highly Reactive Nonheme Iron Intermediate That Oxygenates Strong CH Bonds with Stereoretention. J. Am. Chem. Soc. 2015, 137, 15833−15842. (10) Fan, R.; Serrano-Plana, J.; Oloo, W. N.; Draksharapu, A.; Delgado-Pinar, E.; Company, A.; Martin-Diaconescu, V.; Borrell, M.; Lloret-Fillol, J.; García-España, E.; Guo, Y.; Bominaar, E. L.; Que, L.; Costas, M.; Münck, E. Spectroscopic and DFT Characterization of a Highly Reactive Nonheme Fe V − Oxo Intermediate. J. Am. Chem. Soc. 2018, 140, 3916−3928. (11) Zerk, T. J.; Bernhardt, P. V. Redox-coupled structural changes in copper chemistry: Implications for atom transfer catalysis. Coord. Chem. Rev. 2017, in press. DOI: 10.1016/j.ccr.2017.11.016. (12) Golub, G.; Lashaz, A.; Cohen, H.; Paoletti, P.; Bencini, A.; Valtancoli, B.; Meyerstein, D. The effect of N-methylation of tetraaza-alkane copper complexes on the axial inding of anions. Inorg. Chim. Acta 1997, 255, 111−115. (13) Golub, G.; Cohen, H.; Paoletti, P.; Bencini, A.; Bertini, I.; Messori, L.; Meyerstein, D. Use of Hydrophobic Ligands for the Stabilization of Low-Valent Transition Metal Complexes. 1. The Effect of N-Methylation of Linear Tetraazaalkane Ligands on the Properties of Their Copper Complexes. J. Am. Chem. Soc. 1995, 117, 8353−8361. (14) Nebot-Guinot, A.; Liberato, A.; Máñez, M. A.; Clares, M. P.; Doménech, A.; Pitarch-Jarque, J.; Martínez-Camarena, A.; Basallote, M. G.; García-España, E. Methylation as an effective way to generate SOD-activity in copper complexes of scorpiand-like azamacrocyclic receptors. Inorg. Chim. Acta 2018, 472, 139−148. (15) Acosta-Rueda, L.; Delgado-Pinar, E.; Pitarch-Jarque, J.; Rodríguez, A.; Blasco, S.; González, J.; Basallote, M. G.; GarcíaEspaña, E. Correlation between the Molecular Structure and the Kinetics of Decomposition of Azamacrocyclic Copper (II) Complexes. J. Chem. Soc., Dalton Trans. 2015, 44, 8255−8266. (16) Oberley, L. W.; Spitz, D. R. Handbook of Methods of Oxygen Radicals Research; Grenwald, R. A., Ed.; CRC Press: Boca Raton, FL, 1986. L


Article

Inorganic Chemistry Density. Phys. Rev. B: Condens. Matter Mater. Phys. 1988, 37, 785− 789. (35) Schäfer, A.; Horn, H.; Ahlrichs, R. Fully Optimized Contracted Gaussian-Basis Sets for Atoms Li to Kr. J. Chem. Phys. 1992, 97, 2571−2577. (36) Tomasi, J.; Mennucci, B.; Cammi, R. Quantum Mechanical Continuum Solvation Models. Chem. Rev. 2005, 105, 2999−3093. (37) Cossi, M.; Scalmani, G.; Rega, N.; Barone, V. New Developments in the Polarizable Continuum Model for Quantum Mechanical and Classical Calculations on Molecules in Solution. J. Chem. Phys. 2002, 117, 43−54. (38) Grimme, S.; Antony, J.; Ehrlich, S.; Krieg, H. A Consistent and Accurate Ab Initio Parametrization of Density Functional Dispersion Correction (DFT-D) for the 94 Elements H-Pu. J. Chem. Phys. 2010, 132, 154104. (39) Grimme, S.; Ehrlich, S.; Goerigk, L. Effect of the Damping Function in Dispersion Corrected Density Functional Theory. J. Comput. Chem. 2011, 32, 1456−1465. (40) Schaftenaar, G.; Noordik, J. H. Molden: A Pre- and PostProcessing Program for Molecular and Electronic Structures. J. Comput.-Aided Mol. Des. 2000, 14, 123−134. (41) The PyMOL Molecular Graphics System, Version 1.8; Schrodinger L.L.C.: New York, 2015. (42) Binstead, R. A.; Jung, B.; Zuerbühler, A. D. SPECFIT-32; Spectrum Software Associates: Chapel Hill, NC, 2000. (43) Verdejo, B.; Ferrer, A.; Blasco, S.; Castillo, C. E.; González, J.; Latorre, J.; Máñez, M. A.; Basallote, M. G.; Soriano, C.; GarcíaEspaña, E. Hydrogen and Copper Ion-Induced Molecular Reorganizations in Scorpionand-like Ligands. A Potentiometric, Mechanistic, and Solid-State Study. Inorg. Chem. 2007, 46, 5707−5719. (44) Martell, A. E.; Smith, R. M. Critical Stability Constants; Plenum: New York, 1982; Vol. 5, p 178. (45) Martell, A. E.; Smith, R. M. Critical Stability Constants; Plenum: New York, 1989; Vol. 6, p 214. (46) Bencini, A.; Bianchi, A.; Garcia-España, E.; Micheloni, M.; Ramirez, J. A. Proton Coordination by Polyamine Compounds in Aqueous Solution. Coord. Chem. Rev. 1999, 188, 97−156. (47) Read, R. A.; Margerum, D. W. Kinetics of acid-catalyzed chelate ring opening in nickel(II) polyamine complexes. Inorg. Chem. 1981, 20, 3143−3149. (48) Lan, W.-J.; Chung, C.-S. Kinetics and Mechanism of Dissociation Reactions of Open-Chain and Macrocyclic Tetraamine Complexes of Copper(II). J. Chin. Chem. Soc. 1992, 39, 577−582. (49) Lan, W.-J.; Chung, C.-S. Kinetics of acid-catalyzed dissociation of tetraazamacrocyclic ligand complexes of copper(II). J. Chem. Soc., Dalton Trans. 1994, 0, 191−194. (50) Castillo, C. E.; Máñez, M. A.; Basallote, M. G.; Clares, M. P.; Blasco, S.; García-España, E. Copper(II) Complexes of Quinoline Polyazamacrocyclic Scorpiand-Type Ligands: X-Ray, Equilibrium and Kinetic Studies. J. Chem. Soc., Dalton Trans. 2012, 41, 5617. (51) Soibinet, M.; Gusmeroli, D.; Siegfried, L.; Kaden, T. A.; Palivan, C.; Schweiger, A. Kinetics and Mechanism of the Stepwise Complex Formation of Cu(II) with Tren-Centred Tris-Macrocycles. J. Chem. Soc., Dalton Trans. 2005, 12, 2138−2146. (52) Westerby, B. C.; Juntunen, K. L.; Leggett, G. H.; Pett, V. B.; Koenigbauer, M. J.; Purgett, M. D.; Taschner, M. J.; Ochrymowycz, L. A.; Rorabacher, D. B. Macrocyclic Polyamino Polythiaether Ligands with NxS4-X and NxS5-X Donor Sets: Protonation Constants, Stability Constants, and Kinetics of Complex Formation with the Aquocopper(II) Ion. Inorg. Chem. 1991, 30, 2109−2120. (53) Biver, T.; Lombardi, D.; Secco, F.; Tiné, M. R.; Venturini, M.; Bencini, A.; Bianchi, A.; Valtancoli, B. Kinetic and Equilibrium Studies on the Polyazamacrocycle Neotetren: Metal-Complex Formation and DNA Interaction. J. Chem. Soc., Dalton Trans. 2006, 12, 1524−1533. (54) Biver, T.; Secco, F.; Tinè, M. R.; Venturini, M. Kinetics and Equilibria for the Formation of a New DNA Metal-Intercalator: The Cyclic Polyamine Neotrien/copper(II) Complex. J. Inorg. Biochem. 2004, 98, 33−40.

(55) Siegfried, L.; Kaden, T. A. Formation and Dissociation Kinetics of Cu(II) and Ni(II) Complexes with N2S2-Macrocycles. J. Chem. Soc., Dalton Trans. 2005, 6, 1136−1140. (56) Wilkins, R. G. Kinetics and Mechanism of Reactions of Transition Metal Complexes, 2nd ed.; VCH: Weinheim, Germany, 1991; pp 201− 212. (57) Powell, D. H.; Helm, L.; Merbach, A. E. 17O nuclear magnetic resonance in aqueous solutions of Cu2+: The combined effect of Jahn−Teller inversion and solvent exchange on relaxation rates. J. Chem. Phys. 1991, 95, 9258−9265. (58) Bazzicalupi, C.; Bencini, A.; Bianchi, A.; Ciattini, S.; Giorgi, C.; Paoletti, P.; Valtancoli, B.; Secco, F.; Rosaria Tinè, M.; Venturini, M. Ni(II) and Co(II) complexes with a phenanthroline-containing macrocycle. Thermodynamic, structural and kinetic considerations. Phys. Chem. Chem. Phys. 2000, 2, 4864−4869. (59) Turan, T. S.; Rorabacher, D. B. Steric effects of chelation kinetics. III. Reactions of aquonickel(II) ion with N-alkyl-substituted ethylenediamines. Inorg. Chem. 1972, 11, 288−295. (60) Xu, J.; Huang, W.; McCreery, R. L. Isotope and surface preparation effects on alkaline dioxygen reduction at carbon electrodes. J. Electroanal. Chem. 1996, 410, 235−242. (61) Muscoli, C.; Cuzzocrea, S.; Riley, D. P.; Zweier, J. L.; Thiemermann, C.; Wang, Z.-Q.; Salvemini, D. On the Selectivity of Superoxide Dismutase Mimetics and Its Importance in Pharmacological Studies. Br. J. Pharmacol. 2003, 140, 445−460. (62) Riley, D. P. Functional Mimics of Superoxide Dismutase Enzymes as Therapeutic Agents. Chem. Rev. 1999, 99, 2573−2587. (63) Iranzo, O. Manganese Complexes Displaying Superoxide Dismutase Activity: A Balance between Different Factors. Bioorg. Chem. 2011, 39, 73−87. (64) Ohtsu, H.; Shimazaki, Y.; Odani, A.; Yamauchi, O.; Mori, W.; Itoh, S.; Fukuzumi, S. Synthesis and Characterization of ImidazolateBridged Dinuclear Complexes as Active Site Models of Cu,Zn-SOD. J. Am. Chem. Soc. 2000, 122, 5733−5741.

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