Coordination of Hydrogen Peroxide with Late-Transition-Metal

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

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Coordination of Hydrogen Peroxide with Late-Transition-Metal Sulfonamido Complexes Christian M. Wallen,* John Bacsa, and Christopher C. Scarborough*,† Department of Chemistry, Emory University, 1515 Dickey Drive, Atlanta, Georgia 30322, United States S Supporting Information *

ABSTRACT: Adducts of hydrogen peroxide and transition metals have been implicated as intermediates in biological and industrial processes but have only recently been observed. Therefore, knowledge of how hydrogen peroxide interacts with transition metals is extremely limited. Herein, we report the synthesis of H2O2 complexes of cobalt, nickel, and copper supported by sulfonamido ligands with second-sphere hydrogen bonding. Binding constant and decay kinetics are reported for four new M(H2O2) adducts, providing a foundation for future studies in H2O2 coordination and oxidation catalysis.

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INTRODUCTION Hydrogen peroxide is an attractive oxidant because of its low toxicity, high oxidative efficiency, and benign byproducts.1 Large-scale applications of hydrogen peroxide include the bleaching of wood pulp and textiles and the epoxidation of propylene, processes which involve activation of hydrogen peroxide with metal catalysts.2−4 The nature of the interaction between the metal ion and hydrogen peroxide in these systems and related systems is not thoroughly understood, but a number of chemical species have been proposed as intermediates in computational5−8 and kinetic9−12 studies. Furthermore, M(H2O2) adducts have been computationally implicated to act as secondary reactive intermediates in a variety of cytochrome P450 enzymes.13−18 In a calculated mechanism for cytochrome P450cam enzymes, an FeIII(H2O2) adduct is stabilized by a hydrogen-bond-accepting protein residue and prevented from “uncoupling” (decomplexation of H2O2) by steric pressure from the substrate (A, Figure 1), ultimately forming Compound 1, which is the generally accepted active oxidant in cytochrome P450 enzymes.17,18 In another example, a salophen cobalt complex was shown to aerobically oxidize hydroquinone to p-benzoquinone, and computational studies implicated the key intermediacy of a CoII(H2O2) adduct stabilized by hydrogen bonding to both the hydroquinone substrate and a solvent molecule (B, Figure 1).8 Both of these species react by one-electron reduction of bound H2O2 by the metal center to afford Mn+1−OH hydrogenbonded to a hydroxyl radical. For A, this leads to HAT from the FeIV−OH made accessible by the redox activity of the porphyrin ligand, whereas with B, the CoIII−OH complex is not susceptible to further oxidation, allowing the tethered •OH to react directly with the substrate. Hydroxyl radicals are potent oxidants, and accessing M(H2O2) complexes of redox-active metals would enable such interesting intermediates to be studied experimentally, providing information that could be exploited in the development of novel oxidation catalysts. To © XXXX American Chemical Society

Figure 1. Calculated FeIII(H2O2) intermediate in formation of Compound 1 from Cytochrome P450cam enzyme (A).18 Calculated CoII(H2O2) adduct in aerobic oxidation of hydroquinone by cobalt salophen complexes (B).8 Previously reported M(H2O2) adducts with sulfonamido complexes of zinc19 (C) and cobalt20 (1(H2O2)).

reach this long-term goal, we set out to explore H2O2 coordination to divalent metal centers with electrochemically accessible trivalent oxidation states. Despite multiple theoretical reports,5−18 M(H2O2) adducts were not experimentally observed until 2015, when we published the first unambiguous M(H2O2) adduct, where M is ZnII (C, Figure 1), a report that included lifetime measurements and the first crystal structure with hydrogen peroxide bound to a metal ion.19 In 2016, we reported the first observable M(H2O2) adduct with a redoxactive metal, CoII (1(H2O2), Figure 1), along with lifetime measurements and the first measured binding constant for Received: September 29, 2017

A

DOI: 10.1021/acs.inorgchem.7b02514 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry hydrogen peroxide and a transition metal. 20 While a demonstration of H2O2 coordination to a metal was novel, these initial reports did not elucidate the finer details of this coordination, such as effects of metal identity and ligand electron density on both the lifetime and binding strength of H2O2. Both C and 1(H2O2) are supported by the aryl-sulfonamido ligand [Ts3tren]3−, which stabilizes the otherwise weak M(H2O2) interaction21−24 through second-sphere hydrogen bonding.18−20 We recently reported a related alkyl-sulfonamido ligand, [Bus3tren]3− (Figure 2),25 and demonstrated that the

The preparation of 1P(OH2) and 1P (Scheme 2) has been reported briefly in other contexts;19,20 therefore, the synthesis will be discussed here in depth for clarity. 1H NMR spectroscopic studies of the four-coordinate complexes C and 1 (Figures 1 and 2) indicated a noticeable shift in the proton resonances of the [Bu4N]+ ion upon addition of H2O or H2O2.19,20 This shift is consistent with these axial ligands interfering with ion pairing between the anionic cobalt fragment and the ammonium cation and may also indicate breakup of hydrogen-bonding interactions between these ions in solution. The X-ray crystal structure of 1 reveals that the sulfonyl oxygens are oriented toward the acidic α hydrogens of the [Bu4N]+ counterion, indicating the presence of such hydrogen-bonding interactions in the solid state.25 These solidstate and solution data suggested that the binding of H2O2 to the sulfonamido complexes might be disrupted in part by the [Bu4N]+ counterion. As changing ligand electron density would result in changes in the strength of hydrogen bonding between the anionic cobalt fragment and both H2O2 and an ammonium cation, we turned our attention to identifying a countercation that could not engage in such hydrogen-bonding interactions. We devised the asymmetrically substituted tetraaryl phosphonium ion [tBuArPPh3]+ (Scheme 1) that we hypothesized would Scheme 1. Synthesis of Novel Asymmetric Arylphosphonium Cation

Figure 2. Previously reported19,20,25 sulfonamido ligands and cobalt complexes used in this work.

sulfonyl oxygens of this ligand are significantly more electronrich than those of [Ts3tren]3−. Given the demonstrated importance of intramolecular hydrogen bonding in the coordination of H2O2,18−20 we set out to quantify the relative binding strength of H2O2 to cobalt(II) complexes of [Ts3tren]3− and [Bus3tren]3−. However, crystallographic and spectroscopic studies19,20 revealed hydrogen-bonding interactions between the acidic α protons of the [Bu4N]+ counterion and the ancillary ligand sulfonyl oxygens, an interaction that would complicate interpretation of H2O2 bonding constants. To alleviate this complication, we developed a novel noncoordinating and highly solubilizing tetraaryl phosphonium cation that is much less likely to engage in hydrogen bonding with the ligand sulfonyl oxygens. Furthermore, our work of M(H2O2) adducts thus far has been limited to cobalt and zinc, and in the present study we have expanded to copper and nickel as well. This contribution details our investigation of binding strength and decay kinetics with M(H2O2) adducts of Co, Ni, and Cu, using both aryl- and alkyl-sulfonamido ligands and both ammonium and phosphonium noncoordinating cations. The data on these M(H2O2) adducts establish the properties of M(H2O2) species and give information on the viability of such M(H2O2) intermediates in biological and industrial oxidation mechanisms, thereby illuminating the path to developing novel oxidative methodologies employing H2O2 as the stoichiometric oxidant.

give an improved solubility profile in comparison to [PPh4]+, and the asymmetry of this ion was anticipated to give high crystallinity. In 2008, Charette published a nickel-catalyzed coupling of triphenylphosphine and aryl halides for the preparation of a wide variety of asymmetric tetraaryl phosphonium salts, [ArPPh3]X (X = Cl−, Br−, I−, −OTf).26 Using this procedure, we were able to synthesize [tBuArPPh3][Br] from 1-bromo-3,5-di-tert-butylbenzene and triphenylphosphine (Scheme 1). With this new cation in hand, we were able to synthesize crystalline [ tBu ArPPh 3 ][(Ts 3 tren)Co II (H 2 O)] (1 P (OH 2 ), Scheme 2), which can be dehydrated at room temperature with 3 Å molecular sieves to produce [tBuArPPh3][(Ts3tren)CoII] (1P, Scheme 2). Synthesis of Sulfonamido Complexes of Copper and Nickel. To further explore the effect of redox-active metal ions in the formation of M(H2O2) adducts, we pursued sulfonamido complexes of nickel and copper. For these complexes, the [Bu4N]+ counterion was used, because of their ease of preparation and for direct comparison to C and 1(H2O2). [Ts3tren]3− complexes of copper and nickel were prepared analogously to 1(OH2) in good yields (Scheme 2), and the structures of [nBu4N][(Ts3tren)NiII(OH2)] (3(OH2)) and [nBu4N][(Ts3tren)CuII(OH2)] (4(OH2)) were confirmed by single-crystal X-ray crystallography (Figure 3). The green complex 3(OH2) was dehydrated by heating in THF to 60 °C for several days in the presence of 3 Å molecular sieves, where dehydration was monitored periodically with electronic

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RESULTS Synthesis of [(R3tren)CoII(OH2)]− Complexes with a Hydrogen-Bonding-Innocent Phosphonium Counterion. B

DOI: 10.1021/acs.inorgchem.7b02514 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry Scheme 2. Synthesis of New Metal Sulfonamido Complexesa

a

Yields shown are of pure crystallized products.

Figure 3. Solid-state structures of the anions in 1P(OH2) (left), 3(OH2) (middle), and 4(OH2) (right). Ligand arms have been truncated, and hydrogen atoms on carbon have been omitted for clarity.

absorption spectroscopy. After filtration of the molecular sieves, the orange four-coordinate complex [nBu4N][(Ts3tren)NiII] (3) was crystallized by addition of diethyl ether and pentane. The need for heat to dehydrate this nickel complexes stands in contrast to the ease of dehydration of analogous cobalt complexes, which occurs within a few hours at room temperature with 3 Å molecular sieves.20 Attempts at complete dehydration of the copper complex 4(OH2) were unsuccessful, where only partial dehydration was achieved after several days with molecular sieves at 60 °C. Further attempts to completely dehydrate this complex were unsuccessful. These observations suggest stronger binding of H2O in 3(OH2) and 4(OH2) than in 1(OH2). Binding Constant Measurements. As described above, the nickel and copper complexes (3(OH2) and 4(OH2)) are empirically more difficult to dehydrate than the analogous cobalt complex (1(OH2)). To quantify this observation, we set out to measure the binding constant of water to 3 and 4. Furthermore, we also targeted the measurement of the binding constant of H2O2 to these species to finalize our developing picture of the effects of counterion identity, ligand electronics, and metal identity in H2O2 coordination in this series. Photometric titrations of water into THF solutions of 3 and 4 were performed at room temperature. However, since we have demonstrated that 1(H2O2) is short-lived at room

temperature,20 titrations with H2O2 were performed at −70 °C for 1, 1P, 2, and 3 in THF (Figure 4). For each titration, water or H2O2 was added incrementally to a solution of the four-coordinate complex in THF and the change in electronic absorbance was measured. The titration curve for each complex was modeled mathematically to determine the binding constant (Keq) for each ligand−complex pair, and the resulting Keq values are shown in Table 1. Unfortunately, the difficulty in accessing anhydrous 4 rendered titrations with either water or peroxide impossible under our experimental conditions. Lifetime Measurements. We sought to measure the kinetics for the decay of bound H2O2 to water and oxygen for complexes 1(H2O2), 1P(H2O2), 2(H2O2), and 3(H2O2). In previous work,20 decay of 1(H2O2) could be monitored using 1 H NMR. Unfortunately, this technique could not be used for the related nickel (3) and copper (4) complexes due to paramagnetic broadening that completely obscures the H2O2 signal. Therefore, an alternate method of monitoring H2O2 decay had to be employed. Since the decay of 1(H2O2) to 1(OH2) had been observed to produce O2 gas, we reasoned that the decay of bound H2O2 to H2O would directly correlate to increased pressure in a closed system. Thus, we decided to monitor headspace pressure above a sample of M(H2O2) as a means of measuring decay kinetics. When we monitored the headspace pressure above a sample of 1 and H2O2 in anhydrous C

DOI: 10.1021/acs.inorgchem.7b02514 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry

which can be filtered away. By adding this H2O2/ether solution to anhydrous acetonitrile and then removing the diethyl ether under vacuum, we obtained anhydrous solutions of H2O2 in acetonitrile. Further details of this procedure may be found in the Supporting Information. Decay of M(H2O2) Adducts in Acetonitrile. With the acetonitrile solution of H2O2 in hand, we were able to generate 1(H2O2), 1P(H2O2), 2(H2O2), and 3(H2O2) in acetonitrile by combining 1 equiv of H2O2 with anhydrous 1, 1P, 2, and 3. For each experiment, the four-coordinate complex was stirred in acetonitrile at 25 °C in a sealed system. One equivalent of H2O2 in acetonitrile was added, and the headspace pressure was measured over time using a digital manometer. The pressure data were then fit to an inverse exponential decay function. The quality of the fit in each case demonstrates that decay of bound H2O2 in these complexes follows first-order kinetics. The manometric data and mathematical fits are shown in the Supporting Information, and the calculated half-life values for each complex are shown in Table 2. Since anhydrous 4 is not Table 2. Decay Rates for M(H2O2) Adducta t1/2 (s) 1(H2O2) 1P(H2O2) 2(H2O2) 3(H2O2) 4(H2O2)b

173.5 283.2 66.2 576.4 53.4

± ± ± ± ±

0.6 2.6 0.7 5.2 1.0

R2 0.99959 0.99913 0.99960 0.99606 0.99932

a

Pressure data and mathematical modeling are presented in Supporting Information. bMixture of 4(H2O2) and 4(OH2).

accessible, we estimated the half-life of 4(H2O2) by treating a red solid consisting of an unknown ratio of 4(OH2) and 4 with approximately 1 equiv of hydrogen peroxide and measuring the solution manometrically, which should give an upper estimate of the lifetime of 4(OH2), as protic ligands such as H2O and NH3 increase the lifetime of M(H2O2) adducts through competitive displacement of H2O2.20 Electrochemical Measurements. In an attempt to gain insights into the decay rates of the M(H2O2) adducts, which do not follow any obvious periodic trends, we decided to measure the MII/MIII oxidation potentials for the series of M(OH2) complexes (Figure 5). We considered that there might be a correlation between the MII/MIII oxidation couple and the decay rates of the respective M(H2O2) adducts. Redox potentials for 1(OH2) and 2(OH2) in CH2Cl2 have been reported previously25 but are repeated here for discussion (Table 3). Cyclic voltammetry experiments with electrochemical cycling were performed on 1(OH2), 3(OH2), and 4(OH2) in CH2Cl2 and externally referenced to Fc/Fc+. In the case of 1P(OH2), a phosphonium supporting electrolyte, [tBuArPPh3]PF6, was used instead of [nBu4N]PF6 to avoid any coordinative interactions between the Bu4N+ and sulfonamido complex in solution.

Figure 4. Sample photometric titration data: absorption traces (top) and titration curves with mathematical fits (bottom) for 1 and water at room temperature. Absorption traces and titration curves for 1 with H2O2 and for 1P, 2, and 3 with water and H2O2 are shown in the Supporting Information.

Table 1. Binding Constants for Binding of Water and H2O2 to CoII in 1, 1P, and 2 1 1P 2

H2O

H2O2

2660 ± 360 5760 ± 1260 45100 ± 9230

34.4 ± 1.50 78.7 ± 5.90 17000 ± 4050

THF, we observed an initial increase in pressure followed by a slow decrease in pressure back to the initial pressure. This led us to conclude that the oxygen being produced by disproportionation was being consumed in autoxidation of THF. We sought to avoid this problem by changing the solvent to acetonitrile, which is not susceptible to autoxidation. Preparation of H2O2 Solutions in Acetonitrile. In our previous work, anhydrous solutions of H2O2 were made in THF by extraction from the urea hydrogen peroxide adduct.18,20 Since THF is not compatible with the experimental setup, we needed to find a way to obtain anhydrous H2O2 in acetonitrile. We discovered that anhydrous solutions of H2O2 in diethyl ether were accessible by stirring the urea−hydrogen peroxide adduct in anhydrous diethyl ether and filtering away the solid. Cooling this solution crystallizes any dissolved urea,

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DISCUSSION The following sections will discuss how altering specific variables affects the binding strength of H2O2 to transition metals and the stability of the corresponding M[H2O2] coordination complexes. In summary, the largest effect is observed upon changing the transition-metal ion, although changes in the ligand also show a pronounced effect that highlights the importance of second-sphere hydrogen bonding D

DOI: 10.1021/acs.inorgchem.7b02514 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry

comparison to 1, which bears an aryl-sulfonamido ligand, we expected 2 to have stronger hydrogen-bonding interactions with protic axial ligands, resulting in improved binding of axial protic ligands. As shown above, the use of [Bus3tren]3− on cobalt (2) increases the binding strength for an aqua ligand by a factor of 17 in comparison to the use of [Ts3tren]3− on cobalt (1). Interestingly, the binding strength of H2O2 to cobalt increases by a factor of 500 on moving from [Ts3tren]3− (1) to [Bus3tren]3− (2) (Table 1). Previous studies by Prikhodchenko have shown H2O2 to be a superior hydrogen-bond donor in comparison to water,24,27−29 and our recent work has demonstrated that the O(S) groups of the [Bus3tren]3− ligand are more electron rich than those of [Ts3tren]3−.25 Furthermore, while water is a common ligand in coordination chemistry even in the absence of second-sphere hydrogen bonding, hydrogen peroxide is only known to bind metals when it is assisted by second-sphere hydrogen bonding. From this alone, we anticipated that increasing the hydrogen-bondaccepting character of the ligand would have a more pronounced effect on the binding of H2O2 than on that of H2O, and indeed this is the case. Since the most affordable and practical source of H2O2 is an aqueous solution, the relative binding constants of H2O2 and H2O are a very important factor in the future development of catalysts that utilize coordinated H2O2 intermediates for substrate oxidation. Demonstrating the ability to bias such competitive equilibria is an important result of the present study. Changes in Metal Ion. The titration curves for 3 with water and hydrogen peroxide were too sharp to mathematically calculate a reasonable value for Keq. This is due to a fundamental limitation in measuring binding constants using titration methods. As Keq increases, the mathematically modeled curve relies on an increasingly smaller portion of data, which in turn increases the uncertainty of the mathematical fit. However, the lower limits for Keq can be estimated by comparing simulated titration curves with experimental data (see the Supporting Information). Using this approach, we estimated that the Keq values for the binding of water and H2O2 to NiII in 3 are both above 200000 M−1, much greater than those for the cobalt complex 1 (2660 ± 360 and 34.4 ± 1.50, respectively), consistent with the observation that 3(OH2) is difficult to dehydrate. Unfortunately, dehydration of 4 was not possible under many conditions explored, so that binding constants of H2O and H2O2 to the copper ion in 4 could not be measured. These data highlight that the absolute binding strength of H2O and H2O2 is highly metal dependent. Decay Kinetics of Co(H2O2) Adducts. The half-life of 1(H2O2) in acetonitrile is 173.5 ± 0.6 s, which is shorter than the half-life reported for decay in THF-d8 (353 ± 33 s)20 by a factor of 2. It should be noted that the electronic absorption spectra of 1 in both MeCN and THF are nearly identical and are consistent only with the four-coordinate complex, suggesting that binding of these solvents in the open axial coordination site cannot be detected by this method and that competitive binding of these solvents and H2O2 likely does not account for this difference in half-life. While we have not yet seen evidence of solvent oxidation by M(H2O2) adducts in THF-d8, it is possible that 1(H2O2) has some reactivity with acetonitrile, which could account for the increased rate of decay. The half-life of 2(H2O2) in acetonitrile is 66.2 ± 0.7 s, 2.6 times shorter than that of 1(H2O2). Since Keq for binding of H2O2 to the metal ion in 2(H2O2) is higher than that in

Figure 5. Cyclic voltammograms for 1(OH2) (green), 1P(OH2) (blue), 2(OH2) (purple), 3(OH2) (orange), and 4(OH2) (red). The potential is referenced to Fc/Fc+.

Table 3. Electrochemical Dataa 1(OH2) 1P(OH2)b 2(OH2) 3(OH2) 4(OH2)

Eanodic (mV)

Ecathodic (mV)

E1/2 (mV)

+125 +105 +95 +517 +620

+30 +21 −40 +420d c

+78 +63 +28 +469

a

The scan rate was 100 mV/s with [nBu4N][PF6] as the supporting electrolyte (unless otherwise noted). All values are given vs Fc/Fc+. See Figure 5 for electrochemical traces. Data for 1(OH2) and 2(OH2) have been reported previously and are included here for convenience. b tBu [ ArPPh3][PF6] was used as the supporting electrolyte. cThe return oxidation wave of 4(OH2) is not discernible. dElectrochemically irreversible, but cycling current does not result in changes to the voltammogram (chemically reversible).

in the coordination of hydrogen peroxide to transition-metal ions. Counterion Effects. Replacing the [nBu4N]+ counterion in [(Ts3tren)CoII(L)]− (L = H2O, H2O2) with the phosphonium cation [tBuArPPh3]+ nearly doubles the binding strength for both water and hydrogen peroxide (Table 1). As discussed above, single-crystal X-ray crystallography of 1 and 1H NMR data of 1 and C demonstrate that there is an interaction between the [nBu4N]+ counterion and the sulfonamido ligand oxygen atoms.19,20 Since eliminating cation hydrogen bonding with the cobalt anion increases the binding strength of H2O2 and H2O, it is clear that hydrogen bonding plays a key role in the coordination of these ligands, particularly with H2O2, where coordination in the absence of second-sphere hydrogen bonding has never been experimentally observed. These results suggest that hydrogen bonding of the anionic cobalt species to the countercation results in a reduced ability of the ligand to engage in hydrogen bonding to the axial protic ligand, as the axial ligand and counterion are in competition for interactions with the electron-rich sulfonyl oxygens. As will be seen below, this is the smallest effect on the binding of H2O2 observed in this study. While small, this effect should be considered in the future design of ligands that can support second-sphere hydrogen bonding. Changes in Ancillary Ligand Electronics. Since the alkylsulfonamido complex 2 possesses a more electron rich ligand in E

DOI: 10.1021/acs.inorgchem.7b02514 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry

Summary. Regardless of decay mechanism, the data discussed in this work demonstrate important individual effects of changing ligand electronics, counterion, and metal identity on the lifetime of M(H2O2) adducts and binding strength of hydrogen peroxide. First, the use of the phosphonium cation [tBuArPPh3]+ strengthens the binding of hydrogen peroxide to cobalt sulfonamido complexes. Second, the Keq values for protic ligands binding to the cobalt complex (2) bearing the alkylsulfonamido ligand [Bus3tren]3− demonstrate that increased ligand electron density on the sulfonyl groups strengthens the binding of hydrogen peroxide to a greater extent than the binding of water to the metal ion. Additionally, the nickel complex 3(H2O2) displayed the longest lifetime known for M(H2O2) adducts with redox-active metals and the highest binding constant measured for hydrogen peroxide to a metal ion. Finally, the decreased accessibility of higher metal oxidation states (MIII) with late-transition-metal complexes could be important in avoiding the formation of reactive MIV=O species as would be expected for more reducing metals.30,31 These observations will be useful in developing future catalytic reactions with hydrogen peroxide as the terminal oxidant.

1(H2O2), we initially thought it possible that the faster decay of H2O2 in 2(H2O2) might be a result of a longer residence time of H2O2 on cobalt in 2(H2O2). However, the half-life of 1P(H2O2) (283.2 ± 2.6 s) is 1.6 times longer than that of 1(H2O2) in acetonitrile, even though the former has a higher binding constant. Therefore, we were forced to conclude that there is not a clear connection between the binding strength of H2O2 to 1 and the lifetime of 1(H2O2). Decay Kinetics of Nickel and Copper M(H2O2) Adducts. The nickel complex 3(H2O2), which displayed the highest binding constant, has a half-life of 576.4 s, 3.3 times longer than that of 1(H2O2). The sample containing a mixture of 4(H2O2) and 4(OH2) has the shortest half-life of M(H2O2) complexes examined (