Two Unsupported Terminal Hydroxido Ligands in a Î¼-Oxo-Bridged

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

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Two Unsupported Terminal Hydroxido Ligands in a μ‑Oxo-Bridged Ferric Dimer: Protonation and Kinetic Lability Studies Thomas Philipp Zimmermann,† Thomas Limpke,† Nicole Orth,‡ Alicja Franke,‡ Anja Stammler,† Hartmut Bögge,† Stephan Walleck,† Ivana Ivanovic-Burmazovic,‡ and Thorsten Glaser*,† †

Fakultät für Chemie, Universität Bielefeld, Universitätsstrasse 25, D-33615 Bielefeld, Germany Department Chemie und Pharmazie, Friedrich-Alexander-Universität Erlangen-Nürnberg, Egerlandstrasse 3, D-91058 Erlangen, Germany

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‡

S Supporting Information *

ABSTRACT: The dinuclear complex [(susan){FeIII(OH)(μO)FeIII(OH)}](ClO4)2 (Fe2(OH)2(ClO4)2; susan = 4,7dimethyl-1,1,10,10-tetra(2-pyridylmethyl)-1,4,7,10-tetraazadecane) with two unsupported terminal hydroxido ligands and for comparison the fluorido-substituted complex [(susan){FeIIIF(μ-O)FeIIIF}](ClO4)2 (Fe2F2(ClO4)2) have been synthesized and characterized in the solid state as well in acetonitrile (CH3CN) and water (H2O) solutions. The Fe−OH bonds are strongly modulated by intermolecular hydrogen bonds (1.85 and 1.90 Å). UV−vis−near-IR (NIR) and Mössbauer spectroscopies prove that Fe2F22+ and Fe2(OH)22+ retain their structural integrity in a CH3CN solution. The OH− ligand induces a weaker ligand field than the F− ligand because of stronger π donation. This increased electron donation shifts the potential for the irreversible oxidation by 610 mV cathodically from 1.40 V in Fe2F22+ to 0.79 V versus Fc+/Fc in Fe2(OH)22+. Protonation/deprotonation studies in CH3CN and aqueous solutions of Fe2(OH)22+ provide two reversible acid−base equilibria. UV−vis−NIR, Mössbauer, and cryo electrospray ionization mass spectrometry experiments show conservation of the mono(μ-oxo) bridging motif, while the terminal OH− ligands are protonated to H2O. Titration experiments in aqueous solution at room temperature provide the pKa values as pK1 = 4.9 and pK2 = 6.8. Kinetic studies by temperature- and pressure-dependent 17O NMR spectrometry revealed for the first time the waterexchange parameters [kex298 = (3.9 ± 0.2) × 105 s−1, ΔH⧧ = 39.6 ± 0.2 kJ mol−1, ΔS⧧ = −5.1 ± 1 J mol−1 K−1, and ΔV⧧ = +3.0 ± 0.2 cm3 mol−1] and the underlying Id mechanism for a {FeIII(OH2)(μ-O)FeIII(OH2)} core. The same studies suggest that in solution the monoprotonated {FeIII(OH)(μ-O)FeIII(OH2)} complex has μ-O and μ-O2H3 bridges between the two Fe centers.

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H4hildeMe2, the complex [(hildeMe2){FeIII(μ-O)FeIII}] was isolated. Because of the strong electron donation of the phenolates, the FeIII ions are only five-coordinate.30 Using the slightly less donating carboxylates in H4julia, the FeIII ions are six-coordinate in [(julia){FeIII(OH2)(μ-O)FeIII(OH2)}] with additional terminal water (H2O) ligands to compensate for the diminished charge donation.29 The coordinated H2O ligands can be deprotonated at pK1 = 7.7 ± 0.3 and pK2 = 11.4 ± 0.4. The doubly deprotonated form can be oxidized by O2 in aqueous solution to a transient FeIVFeIII intermediate. On the other hand, the terminal pyridines in susan have π-acceptor character and are thus even less charge-donating than the terminal carboxylates in H4julia. Therefore, the electron demand of the FeIII ions is balanced by anionic terminal ligands such as, e.g., Cl− in [(susan){FeCl(μ-O)FeCl}]2+30 or OAc− in [(susan){Fe(OAc)(μ-O)Fe(OAc)}]2+.33 Here, we show that even terminal hydroxido ligands are present to form the complex [(susan){FeIII(OH)(μ-O)FeIII(OH)}](ClO4)2 (Fe2(OH)2(ClO4)2) with two unsupported terminal FeIII-OH units. For comparison and to exclude a misassignment, we also synthesized the corresponding fluorido complex [(susan)-

INTRODUCTION Terminal FeIII-OH units are proposed to be involved in many biological and biomimetic processes.1−8 Model studies are complicated by the tendency of terminal FeIII-OH units to form thermodynamically stable FeIII(μ-O)FeIII units starting an oligomerization or polymerization process.9 Identification of coordinated OH− ligands by single-crystal X-ray diffraction is not always straightforward. This is due to the fact that OH− cannot be easily differentiated from isoelectronic F−, being a potential source for misassignments of coordinated OH− ligands.10,11 In a rational design, Borovik and co-workers developed a ligand platform based on a triurea derivative of the ligand tren, which stabilizes terminal M-OH and M-O units in the secondary coordination sphere by intramolecular hydrogen bonds and by a sterically protecting cavity.12−18 Nevertheless, complexes with terminal FeIII-OH units have been reported without protecting secondary coordination-sphere effects.19−24 Two terminal hydroxides were even realized in the diferric μoxo complex [(phen)2FeIII(OH)(μ-O)FeIII(OH)(phen)2](NO3)2.25 In our efforts to stabilize and utilize high-valent dimetal cores,26,27 we have developed a bis(tetradentate) dinucleating ligand system with varying terminal donors (Scheme 1).28−33 With the strongly σ- and π-donating terminal phenolates in © XXXX American Chemical Society

Received: July 2, 2018

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DOI: 10.1021/acs.inorgchem.8b01831 Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry

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Scheme 1. Ligand and Complex Abbreviations

Article

EXPERIMENTAL SECTION

Synthesis. The solvents and starting materials were of the highest commercially available purity and were used as received. The ligand susan was synthesized according to the published procedure.30 Caution! Although we experienced no problems, the use of perchlorate salts are potentially hazardous and should only be handled in small quantities and with adequate precautions. [(susan){Fe(OH)(μ-O)Fe(OH)}](ClO4)2·H2O. A solution of susan (116 mg, 0.215 mmol, 1 equiv) in ethanol (EtOH; 12 mL) was added to a solution of Fe(ClO4)2·6H2O (157 mg, 0.433 mmol, 2.01 equiv) in EtOH/H2O (6 mL/2 mL), providing a yellow solution. The addition of triethylamine (NEt3; 0.05 mL, 0.359 mmol, 1.66 equiv) resulted in a burned umber suspension. The suspension was stirred for 1 h, followed by filtration. Diffusion of diethyl ether (Et2O) led to the formation of brownish crystals. Yield: 101 mg (0.112 mmol, 51%). IR (KBr, cm−1): ν̃ 3616 m, 3561 m, 3431 m, 3088 w, 2868 m, 2009 w, 1604 s, 1572 w, 1482 m, 1473 m, 1447 m, 1427 m, 1377 vw, 1366 w, 1353 w, 1309 w, 1294 m, 1275 w, 1093 vs, 1019 m, 1014 m, 1003 m, 974 w, 948 m, 929 w, 863 w, 825 s, 819 s, 777 m, 769 m, 736 w, 723 w, 624 s, 569 m, 551 m, 421 w, 414 w. ESI-MS [methanol (MeOH)]: m/z 364.2 [(susan){Fe(OMe)}2(μ-O)]2+, 341.2 [(susan){FeO}2]2+. ESI-HRMS (MeCN): m/z 799.1725 ({[(susan){Fe(OH)(μ-O)Fe(OH)}](ClO4)}+), 350.1110 ([(susan){Fe(OH)(μ-O)Fe(OH)}]2+). Anal. Calcd for [(susan){Fe(OH)(μ-O)Fe(OH)}](ClO4)2·H2O (C32H46N8Cl2Fe2O12): C, 41.90; H, 5.05; N, 12.21. Found: C, 41.95; H, 5.07; N, 12.14. [(susan){FeF(μ-O)FeF}](ClO4)2·1.8H2O. A solution of susan (209 mg, 0.388 mmol, 1 equiv) in acetone (13 mL) was added to a solution of Fe(ClO4)2·6H2O (290 mg, 0.799 mmol, 2.06 equiv) and (Bu4N)F· 3H2O (306 mg, 0.970 mmol, 2.50 equiv) in EtOH (40 mL), resulting in a greenish-brown solution. After 10 min, NEt3 (0.27 mL, 1.948 mmol, 5.02 equiv) was added, and the reaction was stirred for 2 h, followed by filtration. Diffusion of methyl tert-butyl ether led to the formation of black crystals. Yield: 235 mg (260 mmol, 66%). IR (KBr, cm−1): ν̃ 3577 s, 3462 s, 3103 w, 2881 w, 2019 w, 1603 s, 1571 w, 1476 m, 1447 m, 1429 m, 1377 vw, 1365 w, 1350 w, 1309 w, 1292 m, 1271 w, 1087 vs, 1053 s, 1020 s, 999 m, 979 w, 946 m, 927 w, 864 w, 831 s, 776 m, 735 w, 722 w, 624 s, 540 w, 512 m, 425 w. ESI-HRMS (MeCN): m/z 352.1071 ([(susan){FeF(μ-O)FeF}]2+). Anal. Calcd for [(susan){FeF(μ-O)FeF}](ClO4)2·1.8H2O (C32H45.6N8Cl2F2Fe2O10.8): C, 41.07; H, 4.91; N, 11.97. Found: C, 41.08; H, 4.87; N, 11.85. Crystal Structure Determination. Single crystals of Fe2(OH)2(ClO4)2 and Fe2F2(ClO4)2·H2O were removed from the mother liquor, coated with oil, and immediately cooled to 100(2) K on a four-circle Bruker KAPPA APEX II diffractometer with a 4K CCD detector, Mo Kα radiation, graphite monochromator. Multiscan absorption correction was done with SADABS 2008/1,34 solution and refinement were done with SHELXS/L.35 Crystal data for Fe2(OH)2(ClO4)2 (C32H44Cl2Fe2N8O11; M = 899.35): space group C2/c (No. 15), a = 22.6779(9) Å, b = 15.1175(6) Å, c = 21.9351(8) Å, β = 90.671(2)°, V = 7519.6(5) Å3, Z = 8, T = 100(2) K, μ(Mo Kα) = 0.984 mm−1, ρcalc= 1.589 g cm−3, crystal size = 0.23 × 0.19 × 0.15 mm3, 47409 reflections measured (4.94 ≤ 2θ ≤ 54.00°), and 8195 unique reflections (Rint = 0.0349) used in the refinements. The final R1 values (506 refined parameters) were 0.0359 for 6501 reflections with I > 2σ(I) and 0.0549 for all data. The H atoms of the two OH ligands of Fe2(OH)2(ClO4)2 could be localized and refined with restrained O−H distances. Crystal data for Fe2F2(ClO4)2·H2O (C32H44Cl2F2Fe2N8O10; M = 921.35): space group P3221 (No. 154), a = 11.8915(6) Å, c = 22.8886(14) Å, V = 2803.0(3) Å3, Z = 3, T = 100(2) K, μ(Mo Kα) = 0.997 mm−1, ρcalc= 1.637 g cm−3, crystal size = 0.26 × 0.20 × 0.13 mm3, 39568 reflections measured (6.64 ≤ 2θ ≤ 60.12°), and 5492 unique reflections (Rint = 0.0382) used in the refinements. The final R1 values (287 refined parameters) were 0.0326 for 4824 reflections with I > 2σ(I) and 0.0420 for all data. Absolute structure parameter (Flack) = 0.010(13).36

{FeIIIF(μ-O)FeIIIF}](ClO4)2 (Fe2F2(ClO4)2; Scheme 1). We have examined the protonations of Fe2(OH)22+ in aqueous and acetonitrile (CH3CN or MeCN) solutions and especially determined the structures of the differently protonated species and the kinetics of water exchange. B


Article

Inorganic Chemistry Water-Exchange 17O NMR Measurements. 17O NMR spectra were recorded on a Bruker AVANCE DRX 400WB spectrometer equipped with a spectrospin superconducting wide-bore magnet operating at a resonance frequency of 54.24 MHz at a magnetic induction of 9.4 T. The measurements at atmospheric pressure were performed with a commercial 5 mm Bruker broad-band probe thermostated with a Bruker B-VT 3000 variable-temperature unit. The relaxation rates were measured for both the paramagnetic solutions and metal-free aqueous buffered solutions. The line widths at half-height of the signal were determined by a deconvolution procedure on the real part of the Fourier-transformed spectra with a Lorentzian shape function in the data analysis module of Bruker Topspin 1.3 software. Pressure-dependent measurements were done with a custom-made thermostated high-pressure probe.37 The sample was measured in a standard 5 mm NMR tube cut to a length of 50 mm. To enable pressure transmittance to the solution, the NMR tube was closed with a moveable Macor piston. The advantage of this method is that oxygen-sensitive samples can be easily placed in the NMR tube and sealed with the Macor piston under an argon atmosphere. This facilitates subsequent transfer of the sample to the high-pressure probe. The pressure (2−150 MPa) was applied to the high-pressure probe via a perfluorinated hydrocarbon (hexafluoropropylene oxide, Hostinert 175, Hoechst) and measured by a VDO gauge with an accuracy of ±1%. The temperature (274.2−343.2 K) was adjusted with circulating, thermostated H2O (Colora thermostat WK 16) to ±0.1 K of the desired value and monitored before each measurement with an internal platinum-resistance thermometer with an accuracy of ±0.2 K. 10% enriched 17O-labeled H2O (D-Chem Ltd., Tel Aviv, Israel) was used to give a total enrichment of 1% 17O in the studied samples. Samples were prepared by dissolving solid Fe2(OH)2(ClO4)2 in MeCN and refilling the sample with a degassed buffer (or HClO4) solution to get a 4.8 mM complex solution containing 20% (v/v) MeCN. General Treatment of Water-Exchange 17O NMR Data. The exchange rates of the bound H2O molecules were determined by the line-broadening technique developed by Swift and Connick.38 This approach makes use of the relationship between the reduced transverse relaxation rate (1/T2r) and the mean lifetime of the coordinated solvent (τm) (see eq 1)

(very slow exchange, domain I) was not observed in the available temperature range, and therefore this term can be neglected in the treatment of the data. The dependence of ln(1/T2r) on 1/T clearly shows that under the condition of pH = 2.4 the studied system operates in the slowexchange regime (domain II, where 1/T2r ≅ 1/τm) with significant contribution from 1/T2m at elevated temperatures. In this case, the values of 1/T2r are best described by eq 2, and an exponential Arrhenius-type temperature dependence can be applied in the treatment of the bound H2O relaxation rate 1/T2m, as given by eq 3. 1/T2r = 1/τm{[T2m−2 + (T2mτm)−1 + Δωm 2] /[T2m−1 + τm−1)2 + Δωm 2]}

1/T2m = 1/T2m 0 exp(Em /RT )

(3)

kex = 1/τm = (k b/hT ) exp[(ΔS ⧧/R ) − (ΔH ⧧/RT )]

(4)

The dependence of the exchange rate constant (kex) on the temperature variation can be derived from the Eyring equation (4). Here the reciprocal residence time, or kex, depends on the activation parameters for the water-exchange process, viz., the activation enthalpy, ΔH⧧, and activation entropy, ΔS⧧. A study of the pressure dependence is principally advisible in the slow-exchange domain II because of 1/T2r ≅ kex. The relationship between the applied pressure and exchange rate constant at a fixed temperature T is described by eq 5, kex = kex 0 exp[(−ΔV ⧧/RT )P ]

(5)

where P is the pressure. At pH = 6.2 [0.1 M N,N′-bis(piperazine) (PIPES) buffer, 20% (v/ v) MeCN], nonsignificant line broadening (Δνobs − Δνsolvent < 20 Hz) in comparison to the reference probe was observed (see the Results and Discussion section for the explanation). Cryo Mass Spectrometry (MS) Measurements. Cryospray ionization MS (CSI-MS) measurements were performed on a UHRTOF Bruker Daltonik maXis plus, an electrospray ionization (ESI) quadrupole time-of-flight (qToF) mass spectrometer capable of a resolution of at least 60.000 (fwhm), which was coupled to a Bruker Daltonik Cryospray unit. Detection was in positive-ion mode, and the source voltage was 3.5 kV. The flow rates were 220 μL h−1. Both drying gas (N2), to aid solvent removal, and spray gas were held at 5 °C. The machine was calibrated prior to every experiment via direct infusion of the Agilent ESI-TOF low-concentration tuning mixture, which provided an m/z range of singly charged peaks up to 2700 Da in both ion modes. Spectrophotometric Titration Experiments. UV−vis absorption spectra of Fe2(OH)2(ClO4)2 were recorded in aqueous solution in the pH range 3.8−9.9 by the use of a quartz glass dip-in detector (a light path of 1.0 cm; Hellma Analytics, Müllheim, Germany) coupled to a spectrophotometer (TIDAS S 300, J&M Analytik AG, Essingen, Germany) with flexible light guides. A 20 mL double-wall reaction vessel was used to maintain constant temperature (25 ± 0.1 °C). The desirable pH was achieved by the addition of small amounts of a concentrated sodium hydroxide (NaOH) solution to the solution of 4 × 10−5 M Fe2(OH)2(ClO4)2 in 1.6 × 10−4 M HClO4 (I = 0.1 M, adjusted with NaClO4). After each NaOH addition, the pH was measured using a pH electrode immersed in the Fe2(OH)2(ClO4)2 solution (780 pH-meter, Metrohm). The appropriate correction of the obtained absorption spectra was made as a result of a small dilution effect after the addition of a NaOH solution. Other Physical Measurements. IR spectra (400−4000 cm−1) of solid samples were recorded on a Bruker Vertex 70 as KBr disks. ESIMS spectra were recorded on a Bruker Esquire 3000 ion-trap mass spectrometer equipped with a standard ESI source. UV−vis−NIR absorption spectra were measured on a Shimadzu UV-3101 PC spectrophotometer at ambient temperatures, on a JASCO V770-ST spectrophotometer at ambient temperature, or on an Agilent 8453 diode-array spectrometer coupled to an Unisoko USP-203-A cryostat

π × 1/Pm(Δνobs − Δνsolvent) = 1/T2r = 1/τm{[T2m−2 + (T2mτm)−1 + Δωm 2)/(T2m−1 + τm−1)2 + Δωm 2]} + 1/T2os

(2)

(1)

where T2m describes the transverse relaxation time of coordinated H2O in the inner sphere of the complex in the absence of chemical exchange, ωm is the difference in the resonance frequency of the bulk solvent and solvent in the first coordination sphere, Δνobs − Δνsolvent is the difference between the full line widths at half-height of the 17O NMR signal of the bulk solvent in the presence (Δνobs) and absence (Δνsolvent) of the paramagnetic compound, Pm is the mole fraction of bound H2O {Pm = n(H2O)[complex]/55.56}, and T2os represents an outer-sphere contribution to T2r that arises from long-range interactions of the paramagnetic unpaired electrons of the metal complex with H2O molecules outside the first coordination sphere. The exchange rate constant between the coordinated and bulk solvent, kex, can accordingly be expressed as the reciprocal residence time of the bound solvent molecule kex = 1/τm. The line-broadening experiments were performed at complex concentrations, which assured a reasonable broadening compared to the aqueous reference (Δνobs − Δνsolvent > 20 Hz). The separation of the contributing factors in eq 1 is achieved by measuring the temperature dependence of the reduced transverse relaxation rate (1/T2r). These measurements are, in principle, restricted to a rather small kinetic window between the boiling and freezing points of H2O. For the studied systems, the temperature range from 274.2 to 343.2 K was selected. A contribution of 1/T2os to the reduced transverse relaxation rate that would be clearly visible by a changeover to a positive slope at low temperatures C


Article

Inorganic Chemistry for low-temperature measurements. Cyclic and square-wave voltammograms (CVs and SWs) were measured either by use of an EG&G potentiostat/galvanostat 273A or by use of a VersaStat potentiostat on argon-flushed solutions containing 0.1 or 0.2 M tetrabutylammonium hexafluorophosphate (TBAPF6) as the supporting electrolyte in a conventional electrochemical cell. The working electrode was a glassy carbon (GC) electrode, the counter electrode was a platinum wire, and the reference electrode was Ag/0.01 M AgNO3/CH3CN. The potentials were referenced versus the ferrocenium/ferrocene (Fc+/Fc) couple used as an internal standard. SWs have been recorded with typical frequencies of 15−50 Hz with a GC working electrode. CVs were routinely measured with scan rates of 100 or 200 mV s−1. Magnetic susceptibility data were measured on powdered samples in the temperature range 2−300 K by using a SQUID magnetometer (Quantum Design MPMS XL-7 EC) with a field of 1.0 T. The experimental data were corrected for underlying diamagnetism by use of tabulated Pascal’s constants. The susceptibility data were analyzed on the basis of the usual spin Hamiltonian description for the electronic ground state by using the simulation package JulX written by Eckhard Bill for exchange-coupled systems.39 The spin Hamiltonian employed was Ĥ = − 2JS1̂ ⃗ S2̂ ⃗ +

∑ i = 1,2

gμB Sî ⃗ B⃗

(6)

where J is the exchange coupling constant and g is the average electronic g value. Magnetic moments were obtained from numerically generated derivatives of the eigenvalues of eq 6 and summed up over 16 field orientations along a 16-point Lebedev grid to account for the powder distribution of the sample. 57Fe Mössbauer spectra were recorded on an alternating-constant-acceleration spectrometer. The minimal line width was 0.24 mm s−1 full-width at half-height. The sample temperature was maintained constant in a bath cryostat (Wissel MBBC-HE0106). 57Co/Rh was used as the radiation source. Isomer shifts were determined relative to α-iron at room temperature.

Figure 1. Molecular structures of the cations (a) Fe2(OH)22+ and (b) Fe2F22+. H atoms are omitted for clarity despite those of the coordinated OH− ligands.

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RESULTS AND DISCUSSION Synthesis and Structural Characterization. The reaction of susan with [Fe(OH2)6](ClO4)2 in an EtOH/H2O mixture in the presence of NEt3 and air provided, after diffusion of Et2O into the reaction solution, brownish crystals, which were analyzed by single-crystal X-ray diffraction as Fe2(OH)2(ClO4)2. The molecular structure of Fe2(OH)22+ is shown in Figure 1a (thermal ellipsoid plot in Figure S1a; selected interatomic distances and angles in Table 1). The FTIR spectrum shows two prominent bands at 3616 and 3561 cm−1 assigned to ν(O−H) vibrations of the coordinated OH− as Fe2(OH)2(ClO4)2 crystallizes without solvent molecules. It must be noted that the derivatives of the mononucleating ligand tris(2-pyridylmethyl)amine (tpa), only protonated variants with a {FeIII(μ-O)(μ-O2H3)FeIII} or {FeIII(OH2)(μO)FeIII(OH2)} core but not the {FeIII(OH)(μ-O)FeIII(OH)} core, were crystallographically characterized.40,41 This indicates that the ligand susan is slightly less electron-donating than tpa ligands. Performing the reaction in an EtOH/acetone mixture in the presence of (Bu4N)F·3H2O afforded single crystals of Fe2F2(ClO4)2·H2O (Figures 1b and S1b and Table 1). The high-resolution single-crystal X-ray diffraction data allow one to differentiate between coordinated OH− and F− ligands and to locate the H atoms of the FeIII-OH units from the difference Fourier map in Fe2(OH)2(ClO4)2. ESI-MS on CH3CN solutions confirm the different compositions of Fe2F22+ and Fe2(OH)22+ by m/z 352.11 and 350.11, respectively. In both Fe2(OH)22+ and Fe2F22+, the two terminal ligands (either OH− or F−) coordinate trans to each other with respect to the FeIII(μ-O)FeIII axis. The Fe−OH distances differ

Table 1. Selected Interatomic Distances (Å) and Angles (deg) in Fe2(OH)2(ClO4)2 and Fe2F2(ClO4)2·H2Oa Fe2(OH)2(ClO4)2

Fe2F2(ClO4)2·H2O

Fe1−F1/O1 Fe2−F2/O2 Fe1−O3 Fe2−O3 Fe1−N1 Fe2−N41 Fe1−N2 Fe2−N42 Fe1−N3 Fe2−N43 Fe1−N4 Fe2−N44 Fe1···Fe2

1.9021(16) 1.8479(16) 1.7884(15) 1.8121(15) 2.3051(18) 2.2719(19) 2.2260(18) 2.2573(19) 2.2048(18) 2.1652(19) 2.2597(18) 2.2661(18) 3.5819(4)

1.8520(14)

Fe1−O3−Fe2/Fe1#1

168.33(9)

169.51(13)

1.7884(3) 2.2494(18) 2.2151(17) 2.1501(19) 2.1887(18) 3.5619(6)

Symmetry transformations used to generate equivalent atoms: #1, x − y, −y, −z + 1/3.

a

significantly because of intermolecular hydrogen bonds (Figure 2): Fe1−OH acts as hydrogen-bond acceptor, leading to a longer Fe1−OH bond of 1.90 Å, while Fe2−OH acts as hydrogen-bond donor, leading to a shorter Fe2−OH bond of 1.85 Å. The Fe−OH distance is 1.90 Å in [(phen)2FeIII(OH)(μ-O)FeIII(OH)(phen)2]2+25 and 1.91 Å in [(tpaX)FeIII(μO2H3)FeIII(tpaX)]3+ (tpaX with X = H and 5-Et),40,42 where D


Article

Inorganic Chemistry

Figure 2. Section of the crystal structure of Fe2(OH)2(ClO4)2 to illustrate the intermolecular hydrogen bonding of the coordinated OH− ligands. The OH− of O2 acts as a hydrogen-bond donor, leading to the short Fe2−O2 bond of 1.848 Å. The OH− of O1 acts as a hydrogen-bond acceptor to the coordinated OH− of the neighboring molecule and as a hydrogen-bond donor to a ClO4− anion, leading to a longer Fe1−O1 bond of 1.902 Å. H atoms are omitted for clarity despite those of the coordinated OH− ligands.

the coordinated OH− accepts a hydrogen bond from the coordinated H2O ligand, forming the μ-O2H3 bridging unit. The shorter Fe2−OH bond in Fe2(OH)22+ coincides with almost all other bonds to Fe2, being longer than those to Fe1, which can be explained by charge compensation (even for the bridging oxide: Fe2−μ-O3 1.81 Å; Fe1−μ-O3 1.79 Å). Although the Fe−F bond length of 1.85 Å in Fe2F22+ coincides with the short Fe2−OH, the F− ligands induce significantly shorter Fe1−N4 bonds (2.19 Å) than the corresponding bonds in Fe2(OH)22+ (2.26−2.27 Å), indicating less electrondonating character of the F− ligand. Mössbauer Spectroscopy and Magnetism. The hydrogen-bonding-induced difference in the iron environments is reflected in the Mö ssbauer spectrum (Figure S2a) of Fe2(OH)2(ClO4)2 by two quadrupole doublets with δ1 = 0.45 mm s−1/|ΔEQ|1 = 2.00 mm s−1 and δ2 = 0.44 mm s−1/ |ΔEQ|2 = 1.57 mm s−1. The values for Fe2F2(ClO4)2 are δ = 0.45 mm s−1 and |ΔEQ| = 1.68 mm s−1 (Figure S2b) and thus not distinguishable from the mean values of Fe2(OH)2(ClO4)2. The same holds true for the magnetic properties because both complexes exhibit exchange coupling constants of J = −99 cm−1 (Figure S3). UV−vis−NIR Spectroscopy in CH3CN Solution. The UV−vis−NIR spectrum of Fe2(OH)22+ in CH3CN changes slowly with time, converting to the spectrum of the previously reported [(susan){FeIII(μ-O)(μ-OAc)FeIII}]3+.33 This is in conformity with the hydrolytic activity of μ-oxo-bridged diferric complexes to convert CH3CN to acetamide or acetate.40,43 Hence, UV−vis−NIR spectra were measured directly after dissolution; significant conversion can be ruled out. The spectra of CH3CN solutions of Fe2(OH)22+ and Fe2F22+ show only slight differences (Figure 3a) and are also quite similar to the spectra of [(susan){FeIIICl(μ-O)FeIIICl}]2+30 and [(susan){FeIII(OAc)(μ-O)FeIII(OAc)}]2+.33 Thus, variation of the terminal ligands has no strong influence on the UV−vis−NIR spectra. However, changing the FeIII(μO)FeIII angle by going to the doubly bridged complex

Figure 3. (a) UV−vis−NIR spectra of Fe2(OH)22+ and Fe2F22+ in CH3CN at room temperature and of the free ligand susan for comparison. The dotted lines of the lower inset are solid-state reflectance spectra scaled in intensity to fit the extinction coefficients of the solution spectra. (b) CVs of Fe2(OH)22+ in dimethylacetamide (0.1 M TBAPF6) and Fe2F22+ in CH3CN (0.2 M TBAPF6).

[(susan){FeIII(μ-O)(μ-OAc)FeIII}]3+ induces significant variations in the d−d and charge-transfer (CT) spectral region.33 The absorptions originate mainly from μ-oxo → FeIII ligand-tometal charge-transfer (LMCT) transitions in the range 23000− 35000 cm−1 and pyridine π → π* transitions above 35000 cm−1. Both solution and solid-state reflectance spectra exhibit the characteristic 6A1 → 4T1 d−d transition for μ-oxo-bridged diferric complexes around 10000 cm−1.30,44,45 Only in strongly exchange-coupled complexes can a d−d transition of this high intensity (ε ≈ 5−10 M−1 cm−1) arise for a d5 high-spin ion. Thus, these intense d−d transitions demonstrate the dinuclear μ-oxo-bridged structure of the complexes upon dissolution. Although these observations strongly suggest the structural integrity of Fe2(OH)22+ in a CH3CN solution, the bridging mode may vary by condensation reactions (Scheme 2).43,46 The condensation product of the mono(μ-oxo) core 1 of Fe2(OH)22+ would be the bis(μ-oxo) core 2. This was realized with the ligand tpa6‑Me3 in the complex [(tpa6‑Me3)FeIII(μO)2FeIII(tpa6‑Me3)](ClO4)2,47 which exhibits a broad band around 13200 cm−1 with ε ≈ 80 M−1 cm−1,46 which is absent E


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Inorganic Chemistry Scheme 2. Protonation and Condensation Equilibriaa

a pKa values correspond to the susan complex reported here. Please note that the drawings for the mono(μ-oxo)-bridged complexes only represent the nature of the ligands and not the configurations.

for Fe2(OH)22+. Moreover, the μ-O−FeIII bonds in the diferric complex of core 2 are longer (1.84 and 1.92 Å)47 than those in the core 1 of Fe2(OH)22+ (1.81 and 1.79 Å). Because the intensity of the CT bands reflects the covalency of the metal− ligand bonds,48 the longer μ-O−FeIII bonds in the core 2 result in lower intensities of the μ-oxo → FeIII LMCTs: 31250 cm−1 (4200 M−1 cm−1) and 26700 cm−1 (2000 M−1 cm−1), which are much more intense in Fe2(OH)22+ (Figure 3a).49 This, in conjunction with the similarity of the spectrum of Fe2(OH)22+ to that of Fe2F22+, strongly suggests that Fe2(OH)22+ retains its mono(μ-oxo)-bridged structure 1 in CH3CN solution. The 6A1 → 4T1 transition is at slightly higher energy in Fe2(OH)22+ (9840 cm−1) than in Fe2F22+ (9720 cm−1). Although difficult to quantify, the same trend seems to apply for the 6A1 → 4T2 transition around 17000 cm−1. Because of the negative slopes of these transitions in the Tanabe−Sugano diagram, the higher energies in Fe2(OH)22+ demonstrate a weaker ligand field in Fe2(OH)22+, which could arise from less σ donation and/or more π donation. In Fe2F22+, a sharp spinflip 6A1 → 4A1,4E transition at 21400 cm−1 appears as in the Cl− derivative (20670 cm−1).30 Although this transition is less well resolved in Fe2(OH)22+, it shifts to lower energy (19000− 20000 cm−1). This transition is independent of the ligand field and represents the strength of d−d electronic repulsion. The lower value for Fe2(OH)22+ indicates a higher covalency, making a stronger π donation of OH− versus F− more likely for the explanation of the weaker ligand field. This is in accordance with the structural analysis. Electrochemical Characterization. In contrast to these small differences for Fe2(OH)22+ and Fe2F22+ in the magnetic properties and the UV−vis−NIR and Mössbauer spectra, the electrochemical properties differ significantly (Figure 3b). Fe2F22+ exhibits an irreversible oxidation at 1.40 V versus Fc+/ Fc, which shifts to 0.79 V versus Fc+/Fc in Fe2(OH)22+. This 0.61 V (!) cathodic shift demonstrates a much stronger electron donation of the OH− ligands in Fe2(OH)22+ than of the F− ligands in Fe2F22+. The stronger electron donation was already evident in the structural characterization because the OH− ligand induces longer Fe1−N4 bonds than the F− ligand (vide supra).

In order to obtain some insight into this strong cathodic shift, we compare the potentials found in other complexes of this isostructural series: 1.48 V versus Fc+/Fc in [(susan){FeIIICl(μ-O)FeIIICl}]2+30 and 1.45 V versus Fc+/Fc in [(susan){FeIII(OAc)(μ-O)FeIII(OAc)}]2+.28 Thus, not the potential of Fe2F22+ but that of Fe2(OH)22+ is exceptional. Interestingly, we have also obtained the disordered complex [(susan){FeIIICl(μ-O)FeIII(Cl)0.75(OCH3)0.25}]2+, where the species with methanolate ligation exhibits a lower potential for oxidation at 1.14 V versus Fc+/Fc.30 The UV−vis−NIR spectra provided a lower ligand field in Fe2(OH)22+ than in Fe2F22+ (vide supra). A lower ligand field arises from less σ donation and/or more π donation. Because the electrochemical and structural characterization indicate that the OH− ligand is more charge-donating than the F− ligand, the lower ligand field indicates that there is a strong contribution of higher π donation in the overall stronger electron donation. This is not surprising because hydroxide OH− ligands are the simplest homologues of alcoholate OR− and phenolate OPh− ligands that are both strong π donors. Titration Experiments in CH3CN Solution Observed by UV−vis−NIR and Mössbauer Spectroscopies. Protonation of the coordinated OH− ligands was investigated in CH3CN solutions of Fe2(OH)22+ at −40 °C. In the UV−vis spectra, the addition of 1 or 2 equiv of HClO4 results in shifts in the μ-oxo → FeIII LMCT and d−d regions (Figure 4a). The strong μ-oxo → FeIII LMCT at 28820 cm−1 in Fe2(OH)22+ shifts to lower energies (first protonation, 27300 cm−1; second protonation, 26800 cm−1) and retains a strong intensity, in agreement with the persistence of a monobridged {FeIII(μO)FeIII} core in all of the protonation states, i.e., the formation of cores 3 and 5 and not their condensation products 4 and 6. At lower energies, two protons lead to a shift of the 6A1 → 4T2 transition from 17850 to 16000 cm−1, demonstrating a stronger ligand field due to the formation of two less πdonating H2O ligands. Interestingly, the monoprotonated species seems to have both transitions in this region in conjunction with one OH− and one H2O ligand. The transitions of the monoprotonated species closely resemble those of [(tpaR)FeIII(OH)(μ-O)FeIII(OH2)(tpaR)]3+ complexes (R = H) of core 3.40,42 The condensed form 4, F


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Inorganic Chemistry

experiments. Single (Figure S4b) and double (Figure S4c) protonations provide no significant changes, indicating the integrity of the overall molecular structure in conjunction with the formation of 3 and 5. Spectrophotometric Titration in Aqueous Solution. To establish the nature of the diiron species present in an aqueous solution as a function of the pH, a spectrophotometric titration in the pH range 3.8−8.5 was performed for Fe2(OH)22+ (Figure S5). The UV−vis spectral changes show the existence of two sets of isosbestic points (inset of Figure S5), indicating the presence of more than two species in the solution in the studied pH range. The corresponding plots of absorbance measured at four different wavenumbers (20000, 27250, 28570, and 29000 cm−1) versus pH (Figure S6) clearly show the existence of two acid−base equilibria for which pKa values were determined using pKa-based equations according to Polster and Lachmann.52 The two pKa values, pK1 and pK2, obtained at various wavelengths are summarized in Table S1 and result in the mean values of 4.9 and 6.8, respectively. The acid−base equilibria related to the determined pKa values are defined in Scheme 2. The UV−vis spectra of the three distinct species 1, 3, and 5 of the μ-oxo-bridged diiron complex generated at specified pH values in aqueous solution are presented in Figure 5.

Figure 4. (a) Protonation of Fe2(OH)22+ in CH3CN at −40 °C with 1 and 2 equiv of HClO4, followed by UV−vis−NIR spectroscopy. The experiments were performed with concentrations of 0.83 mM (left axis) and 0.23 mM (right axis). (b) Time traces at selected wavenumbers during consecutive protonation and deprotonation reactions. The experiment was performed on a 0.83 mM CH3CN solution of Fe2(OH)22+ at −40 °C.

Figure 5. UV−vis spectra of the three distinct species 1, 3, and 5 formed in aqueous solution of the μ-oxo diiron complex Fe2(OH)22+ at the pH provided.

realized in [(tpaR)FeIII(μ-O)(μ-OH)FeIII(tpaR)]3+ (R = 6-Me), exhibits a characteristic transition at 18200 cm−1 with a strong intensity (ε ∼ 700−800 M−1 cm−1).46 This has also been observed with other ligands.50,51 Such a transition is definitively absent in the monoprotonated species studied here. An additional argument against the condensed form 6 for the doubly protonated species is the strong μ-oxo → FeIII LMCT, which would be absent in the bis(μ-hydroxido) core 6. Thus, the protonations of Fe2(OH)22+ in CH3CN can be assigned to the reactions from 1 to 3 and from 3 to 5. These protonations seem to be almost reversible by consecutive additions of HClO4 and NEt3 (Figure 4b). Analogous protonation experiments were performed on more concentrated solutions for Mössbauer spectroscopy (Figure S4). Dissolution of Fe2(OH)22+ in CH3CN (Figure S4a) provides no significant changes to the Mössbauer spectrum in the solid (Figure S2a). This confirms the assignment to 1 in CH3CN solution by the UV−vis−NIR

The spectra of these three protonated species in aqueous solution differ slightly from those in CH3CN, which can be ascribed to different kinds of solvation effects. However, the overall spectral features are conserved. The spectrum (Figure 5) of Fe2(OH)22+ [31800 (sh), 28300, and 17400 cm−1] also does not resemble the characteristic transitions for a complex of core 2 at 13200 and 21300 cm−1 (vide supra). This, together with the rather high extinction bands in the UV region, suggests that the mono(μ-oxo)-bridged core 1, and not the product of its condensation 2, exists as the main species in aqueous solution at higher pH. The spectrum of the monoprotonated species 3 [31700 (sh), 29000 (sh), 20100 and 16800 (sh) cm−1] resembles the electronic spectrum of [(tpa)Fe(OH)(μ-O)Fe(OH2)(tpa)]3+, containing a hydrogenbond-bridged (μ-O2H3)− moiety.40 In analogy to the CH3CN solution, the formulation of the monoprotonated species as the G


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Inorganic Chemistry

Scheme 3. Reversible Conversion from Trans to Cis Configuration by Going from a Monobridged to a Dibridged Diferric susan Complex Observed Recently in a Carboxylate Shift33

Figure 6. (A) Plot of ln(1/T2r) versus 1/T to determine the water-exchange rate constant and activation parameters for the dimer complex at pH = 2.4. (B) Plot of ln(kex308) versus pressure at pH = 3.0 and 308 K. Experimental conditions: [Fe2(OH)22+] = 4.8 × 10−3 M, I = 0.1 M (pH and ionic strength adjusted with HClO4 and NaClO4, respectively), in the presence of 20% (v/v) CH3CN and of 10% enriched 17OH2 to give a total enrichment of 1% 17O in the studied samples.

exchangeable site for phosphodiester binding and a nucleophilic hydroxido group on the adjacent Fe site. ESI-MS Experiments in Aqueous Solution at Different pH Values. The nature of the species formed in acid−base equilibria of Fe2(OH)22+ has been further explored by ESI-MS experiments. On the basis of ESI-MS experiments performed in H2O, the diiron complex mainly exists in its fully protonated form 5 at pH = 2.5 (pH adjusted with HClO4). Although we could not directly see the species with two coordinated H2O molecules, the main species detected under these conditions are m/z 368.0769 and 400.0666 (Figure S7), which correspond to [(susan){Fe(μ-O)Fe}]4+ as adducts with different anions being present (Cl− and ClO4−). Because water ligands are very labile and easily released under applied ESI-MS experimental conditions, the detected species are indicative for the core 5. Upon going to higher pH, peaks ascribed to the complex with one coordinated hydroxide ligand, i.e., [(susan)FeIII(OH)(μ-O)FeIII]3+ (see species with m/z 359.094 and 391.084 in Figure S8), become more prominent than those in the spectrum obtained at pH = 2.5. This species is related to the monoprotonated form 3 (where one H2O molecule does not remain coordinated under MS conditions) or its condensation product 4. However, on the basis of the UV− vis characteristics, the latter is ruled out. In the ESI-MS spectrum obtained at even higher pH = 6, the peaks related to species 5 disappear and those related to species 3 become less intense, whereas peaks described to doubly bridged complex 2, [(susan){FeIII(μ-O)2FeIII}]2+ (m/z 341.106), and the parent

condensed core 4 can be excluded by the absence of the characteristic band at 18200 cm−1 43,46 during pH titration in aqueous solution. This is consistent with the huge excess of H2O molecules, which should suppress the condensation reaction even more than in CH3CN solution. It is important to note that the existence of a second bridge in the form μ-O2H3 cannot be excluded. It should be noted that the formation of a second bridge requires that the two coordination sites for the exogenous ligands must change from trans to cis configuration. This transformation requires also a movement of the donor atoms of the ligand susan. Although such a transformation might be thought to be energetically too demanding, we have observed such a transformation recently in a carboxylate shift (Scheme 3).33 Thus, such a transformation is also feasible for these protonation/deprotonation equilibria. For the doubly protonated form, the spectrum with its broad and weak bands between 14300 and 16700 nm and the persistence of the strong LMCT bands also indicate that this is the monobridged diiron species 5 with two terminal H2O ligands and not its condensation product 6. The acid−base titration experiments showed that the studied complex behaves as a diacid in aqueous solution. Notably, the presence of two successive titrable protons with pKa values of 5.00 and 6.85 has been demonstrated in potentiometric titration of the complex [(phen)2Fe(OH2)(μO)Fe(OH2)(phen)2](NO3)4.25 Importantly, its monoprotonated form analogous to 3 formed under the conditions of ca. pH = 6 was found to be the catalytically active species during hydrolysis of phosphodiesters because it provided both an H


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Inorganic Chemistry

sensitive to the entropy than to the volume).55,56 The dissociative character of the water exchange could be an indication of the labilizing effect of the μ-oxo group, as was revealed in the case of μ-oxo-bridged chromium, ruthenium, and rhodium complexes possessing μ-oxo groups in the cis or trans position to the leaving group.57−59 However, in view of the absence of water-exchange data for the corresponding mononuclear iron(III) complexes bearing similar ligands (for example, the mononuclear iron complexes of tpa), it is difficult to determine the real effect of the μ-oxo group in the cis position on the rate and mechanism of water exchange at the Fe centers of 5. It can be stated that the water exchange at each FeIII center of the μ-oxo dimer proceeds much faster and more dissociatively than that at the [FeIII(OH2)6]3+ ion (kex298 = 1.6 × 102 s−1; ΔV⧧ = −5.4 ± 0.4 cm3 mol−1).60,61 Interestingly, the rate of water exchange at the FeIII centers of 5 is on the same order of magnitude as that measured for [FeIII(H2O)5(OH)]2+, with the hydroxido group in the trans position to the exchanging H2O molecule (kex298 = 1.4 × 105 s−1; ΔV⧧ = ± 7.0 + 0.3 cm3 mol−1.60,61 This demonstrates that the net labilizing effect of the ligand susan and the μ-oxo group in the cis position to the leaving H2O molecule in 5 is similar to the labilizing effect of the trans-hydroxido group in [FeIII(H2O)5(OH)]2+, leading in both cases to the interchange dissociative mechanism. A similar Id mechanism has been proposed for the water-exchange reactions at the FeIII centers of polyaminecarboxylate complexes, which display waterexchange rate constants 1−3 orders of magnitude higher than those of 5. This can be ascribed partially to the labilizing influence of the polyaminecarboxylate chelate and partially to the change in the coordination number from 6 to 7 in such chelate complexes.62 Water-exchange experiments performed for the μ-oxobridged diiron complex in a buffered solution at pH = 6.2 where 3, i.e., {FeIII(OH)(μ-O)FeIII(OH2)}, should exist as the main species did not display any significant line broadening (Hz) in comparison to the reference (Figure 7), indicating that either no water exchange at the Fe centers can be observed in the time window of the applied 17O NMR technique or there is no H2O coordinated to the Fe centers.

complex 1, i.e., [(susan){FeIII(OH)(μ-O)FeIII(OH)}]2+ (m/z 350.1115), appear as the main peaks (Figure S9). The observed trends in the ESI-MS experiments upon pH changes are in accordance with the proposed acid−base equilibria and complex speciation in aqueous solutions (Scheme 2). However, it should be noted that, under conditions of higher pH, the peak detected at m/z 341.106 indicates the existence of the condensation product of the dihydroxo form 1, i.e., the doubly bridged form 2 that was not observed in UV−vis spectra during pH titration. One possible explanation for the occurrence of the condensation products in ESI-MS experiments may involve the fact that, during evaporation of the H2O from the probe (even H2O bound to the metal center does not remain intact), the equilibrium between 1 and 2, which is very sensitive to the amount of H2O present, is shifted toward formation of the condensation products. Temperature- and Pressure-Dependent 17O NMR Studies on Water Exchange at the Fe Center of the μOxo Dimer Complex. In principle, water exchange at the Fe center of Fe2(OH)22+ should be observed under conditions of low and near-neutral pH, i.e., at pH values of about 2−3 and 5−6, where the main species present in the solution are 5 and 3, respectively. Water-exchange experiments performed at pH = 2.4 revealed small but significant line broadening, indicating the occurrence of water exchange at the FeIII center. The plot of ln(1/T2r) versus 1/T (Figure 6a) enabled one to determine the water-exchange rate constant and activation parameters for the studied complex under such conditions. The obtained results show that water exchange at each FeIII center of the dinuclear complex at pH = 2.4 proceeds with kex298 = (3.9 ± 0.2) × 105 s−1 and the activation enthalpy and entropy equal to ΔH⧧ = 39.6 ± 0.2 kJ mol−1 and ΔS⧧ = −5.1 ± 1 J mol−1 K−1, respectively. Pressure-dependent H217O NMR experiments performed at pH = 3 and 308 K in the pressure range 2−150 MPa (Figure 6b) revealed the activation volume equal to ΔV⧧ = +3.0 ± 0.2 cm3 mol−1. The value of kex298 ∼ 105 s−1 for the water exchange for 5 is in a typical range for FeIII centers known from the literature.53 These values were shown to be strongly affected by the electronic and structural properties as well as the charge of the coordinated ligand, displaying kex298 ∼ 104 s−1 for more positively charged ligands and kex298 ∼ 107 s−1 for more negatively charged ligands.53,54 Although we suppose that the presence of the μ-oxo group should show a labilizing effect on the coordinated water, we did not find any example in the literature where water exchange at a {FeIII(OH2)(μ-O)FeIII(OH2)} core was measured and thus direct comparisons are not possible. The lack of literature data can result from the fact that the signals obtained by the 17O NMR line-broadening technique for such systems are rather small because of antiferromagnetic coupling between the two high-spin FeIII centers. The relatively small value of the activation entropy together with the small positive value of the activation volume are an indication for the dissociatively activated interchange mechanism (Id) in which breaking of the FeIII−OH2 bond dominates the whole process of water exchange at the FeIII center by a small margin only. The slightly negative entropy of activation suggests that, during the activation process with a dissociative interchange mode, two interchanging H2O molecules may interact with the O atom of the μ-oxo group via hydrogen bonding (the interaction in the second coordination sphere such as hydrogen bonding as well as dipole interaction is more

Figure 7. Nonsignificant line broadening (Hz) at pH = 6.2 ([Fe2(OH)22+] = 4.8 × 10−3 M, 0.1 M PIPES buffer, 20% CH3CN, 1% 17OH2) in comparison to the reference probe studied by the 17 OH2 NMR technique. I


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Inorganic Chemistry

complex {FeIII(OH)(μ-O)FeIII(OH)}, the monoprotonated species {FeIII(OH)(μ-O)FeIII(OH2)}, and the doubly protonated species {FeIII(OH2)(μ-O)FeIII(OH2)}. In all cases, the presence of corresponding condensation products could not be confirmed. Interestingly, the absence of an exchangeable H2O molecule in the monoprotonated species {FeIII(OH)(μO)FeIII(OH2)} suggests a strong hydrogen-bonding interaction between the OH− and H2O ligands and, hence, the existence of a (μ-O2H3) bridge between two FeIII centers in addition to the μ-O bridge. We have investigated for the first time by 17O NMR spectrometry the kinetic lability of aqua ligands within a {FeIII(OH2)(μ-O)FeIII(OH2)} core, demonstrating that it correlates with the lability of the simple [FeIII(H2O)5(OH)]2+. Thus, it seems that the labilizing effect of the μ-oxo group in the cis position to the leaving H2O molecule is similar to the labilizing effect of the transhydroxido group. This, as well as our overall results, provides important information in the field because iron(III) hydroxido and iron(III) oxo motifs play a role whenever iron species, in molecular form or within materials, are involved in stoichiometric or catalytic reactions under aerobic conditions in aqueous solutions and even on surfaces.

Because the latter case can be excluded on the basis of the pH titration and UV−vis−NIR experiments, the obtained result, although it appears to be inconsistent on first sight, can be explained in view of the fact that the H2O molecule in the monoprotonated complex at pH = 6.2 is hydrogen-bonded to the hydroxido ligand to form a (μ-O2H3)− bridge between the two Fe centers. It is supposed that the exchange of such bridged H2O molecules must be very slow and therefore cannot be observed in the time window of the applied 17O NMR technique. Additionally, the antiferromagnetic coupling between two high-spin FeIII centers in such a (μ-O)(μ-O2H3)bridged species could be even stronger and can result in the even weaker signals obtained by the 17O NMR technique than in the case of the mono(μ-oxo)-bridged complex 5 for which the line broadening was already small but significant.

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CONCLUSION By independent synthesis and thorough solid-state and solution characterization of Fe 2 (OH) 2 (ClO 4 ) 2 and Fe2F2(ClO4)2, we could clearly distinguish them and identify their similarities/differences. Because iron complexes with terminal FeIII-OH units are involved in a variety of biological and chemical processes, this is of a significant importance because coordination of OH− can easily be mistaken with coordination of isoelectronic F−. The most prominent difference between these two complexes is reflected in their oxidation potentials, where the increased electron donation by OH− results in a redox potential that is lower by 610 mV than in the case of F− coordination. Analysis of the structural parameters and d−d transitions identifies a much stronger covalency and π donation of the OH− ligand than the F− ligand as the main contribution. The solid-state structure of Fe2(OH)2(ClO4)2 shows a strong variation of the Fe−OH bond lengths modulated by intermolecular hydrogen bonds. When the coordinated OH− acts as a hydrogen-bond donor, the Fe−OH bond is short (1.85 Å); when the coordinated OH− acts as a hydrogen-bond acceptor, the Fe−OH bond is long (1.90 Å). Thus, hydrogen bonding can strongly modulate the Fe−O covalency. This influence of hydrogen bonding on the metal−ligand covalency and hence electronic structure reflected, e.g., in the variation of reduction potentials, has already been proposed for model complexes and proteins.63−65 The strong variation of the Fe− OH bond lengths by hydrogen bonds, detected here for Fe2(OH)22+, provides a direct observation of this effect. The complex Fe2(OH)22+ has been studied intensively in CH3CN and aqueous solutions by several spectroscopic methods. Especially, protonation/deprotonation was studied and provided two reversible acid−base equilibria equilibria in both solvents. Titration experiments in aqueous solution at room temperature provide the pKa values as pK1 = 4.9 and pK2 = 6.8. Comparing these values to pK1 = 7.7 and pK2 = 11.4 determined for [(julia){FeIII(OH2)(μ-O)FeIII(OH2 )}],29 which exhibits a terminal carboxylato donor instead of a pyridine donor, shows that the hydroxido ligands in Fe2(OH)22+ are more difficult to protonate. This stronger preference for the more charge-donating hydoxido ligands indicates a lower electron density of the FeIII ions in Fe2(OH)22+. Thus, the carboxylato donors of the ligand (julia)4− are more charge-donating than the pyridine donors of the ligand susan. Our UV−vis−NIR, Mössbauer, titration, cryo-ESI-MS, and water-exchange studies have revealed the nature of the parent

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.8b01831. pKa values, thermal ellipsoid plots, Mössbauer spectra, magnetic data, titrations followed by UV−vis spectroscopy and analysis, and UHR-ESI-MS spectra (PDF) Accession Codes

CCDC 1815753−1815754 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.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected] (T.G.). ORCID

Ivana Ivanovic-Burmazovic: 0000-0002-1651-3359 Thorsten Glaser: 0000-0003-2056-7701 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We thank Bielefeld University for financial support. REFERENCES

(1) Sträter, N.; Klabunde, T.; Tucker, P.; Witzel, H.; Krebs, B. Crystal structure of a purple acid phosphatase containing a dinuclear Fe(III)-Zn(II) active site. Science 1995, 268, 1489−1492. (2) Lah, M. S.; Dixon, M. M.; Pattridge, K. A.; Stallings, W. C.; Fee, J. A.; Ludwig, M. L. Structure-function in Escherichia coli iron superoxide dismutase: Comparisons with the manganese enzyme from Thermus thermophilus. Biochemistry 1995, 34, 1646−1660. (3) Sträter, N.; Lipscomb, W. N.; Klabunde, T.; Krebs, B. TwoMetal Ion Catalysis in Enzymatic Acyl- and Phosphoryl-Transfer Reactions. Angew. Chem., Int. Ed. Engl. 1996, 35, 2024−2055.

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