Analysis of the Active Species Responsible for Water Oxidation Using a Pentanuclear Fe Complex.

Water splitting with sunlight is today one of the most promising strategies that can be used to start the imperatively needed transition from fossil to solar fuels. To achieve this, one of the key reactions that need to be mastered is the electrocatalytic oxidation of water to dioxygen. Great developments have been achieved using transition metal complexes mainly based on Ru, but for technological applications it is highly desirable to be able to use earth-abundant transition metals. The intrinsic chemistry of first row transition metals and in particular the lability of their M-L bonds in water imposes serious challenges for the latter to work as real molecular catalysts. The present work addresses this issue based on a molecular pentanuclear Fe5 complex and describes the different protocols and tests that need to be carried out in order to identify the real active species, responsible for the generation of dioxygen.

Introduction

One of the main threats on the lifestyle of our modern societies is the increased global warming effect caused through the emission of greenhouse gases. The massive burning of fossil fuels over the course of the past decades has resulted in an alarming increase in carbon dioxide concentrations responsible for the global climate change and concomitant environmental issues.

It is thus extremely urgent to replace fossil fuels by new energy conversion schemes based on clean and environmentally respectful fuels. A potential option is the production of H2 through water splitting with sunlight (hν-WS) as indicated in Equation (1).

Hydrogen generated in this manner is termed solar fuel. Furthermore, hν-WS is also termed as artificial photosynthesis (Grätzel, 1981; Berardi et al., 2014; Lewis, 2016; Nocera, 2017; Roger et al., 2017; Guan et al., 2018) because there are a number of analogies with this reaction and the one that occurs in the natural photosynthesis summarized in Equation (4) (Nelson and Ben-Shem, 2004; McEvoy and Brudvig, 2006; Croce and van Amerongen, 2014).

From a chemical perspective these analogies include:

(1) Both reactions are thermodynamically uphill, driven by sunlight and require the participation of light harvesting agents that can transfer the sun's energy into high-energy-density chemicals such as carbohydrates or H2; (2) both processes need catalysts to speed up the redox reactions; and (3) the water oxidation reaction occurs in an identical fashion in natural photosynthesis and in hν-WS. In the former case a tetramanganese cluster located in photosystem II (PSII) is employed as a catalyst to speed up the water to dioxygen reaction.

The water oxidation reaction has long been regarded as a major bottleneck that ought to be solved to be able to develop devices based on hν-WS. However, during the last decade, a large degree of knowledge has been generated based on both oxide materials (Smith et al., 2013; McCrory et al., 2015; Godwin et al., 2018) and molecular transition metal complexes (Blakemore et al., 2015; Garrido-Barros et al., 2017; Matheu et al., 2019a, 2019b).

In this context, the molecular water oxidation catalysis field has experienced a significant progress over the last 10 years powered by the promise of generating sustainable carbon neutral fuels based on water splitting (Lewis and Nocera, 2006, 2007; Llobet and Meyer, 2011). A particularly noteworthy contribution has been the development of water oxidation catalysts containing molecular Ru complexes owing to the high degree of understanding of their performance at a molecular level. This has been achieved thanks to a thorough description of their reactivity, electrochemical and spectroscopic properties, as well as a detailed characterization of their reaction intermediates, all combined with complementary computational models (Ellis et al., 2010; Radaram et al., 2011; Sala et al., 2014; Matheu et al., 2015; Keidel et al., 2017).

It would be very convenient to use non-toxic earth-abundant transition metal complexes such as Fe-based catalysts as water oxidation catalyst (WOCs) for the generation of technologically useful devices. Few Fe-based WOCs have been reported so far, but most of them are unfortunately not free from controversy since in most of the cases the real active catalytic species is most likely the corresponding oxide rather than the initial molecular complex (Hoffert et al., 2013; Pattanayak et al., 2017). Therefore, rigorous analysis of the active species is essential in order to achieve meaningful information of the catalytic species, not only in water oxidation but in the field of molecular electrocatalysis (Kaeffer et al., 2016; Folkman et al., 2018).

The present work analyzes the water oxidation catalysis initially associated to the complex [FeII4FeIII(μ3-O) (μ-bpp)6]3+, from now on labeled as [FeFe]3+ or Fe3+ (Okamura et al., 2016), where bpp− is the anionic tetradentate dinucleating bridging ligand 3,5-bis(2-pyridyl)pyrazolato; see Figure 1 for a drawing of its two electron oxidized derivative. The transformation of the metal complex into the corresponding oxide is a major concern here since it precludes the correlation of reactivity with the catalytic process and thus becomes a futile exercise. Such exercises could generate misleading information in the water oxidation field.

Figure 1

Crystal Structure of Fe55+

Left, ball and stick drawing of the cationic part of [FeFe]5+. The Fe centers and the oxygen atoms are represented as green and red spheres, respectively. The two Fe(II) centers are situated in the equatorial plane of the bipyramid. The bpp− ligand is represented with gray sticks. H atoms are not drawn for simplicity reasons. Top right, partial representation of the upper part of bipyramidal structure of the Fe5+ complex showing only two of the six bpp− ligands (for simplicity purposes) bridging axial and equatorial Fe centers. Bottom right, drawn structure of the bpp− (L−) ligand and its representation using arcs connected with N.

This manuscript summarizes the main critical tests that need to be carried out to prove the molecular nature of the catalytic processes using Fe3+ but that it obviously can be extended to other iron complexes as well as to other earth-abundant first row transition metal complexes.

Results

Fe5n+ Synthesis, Structure, and Redox Electronic Properties in MeCN

The reaction of Fe(SO4)∙7H2O and bpp− in MeOH in an open atmosphere gives rise to the formation of a pentanuclear complex [FeII4FeIII(μ3-O) (μ-bpp)6]3+, [FeFe]3+ or Fe3+, with a relatively high isolated yield of ∼72%. Although UV-vis spectroscopy shows that the reactions proceed quantitatively (see Figure S3), the decrease in the isolated yield is due to the crystallization process and the follow-up operations needed to isolate a pure solid.

Complex Fe3+ can be cleanly and successively oxidized by two one-electron processes in MeCN reaching the high oxidation state complex [FeII2FeIII3(μ3-O) (μ-bpp)6]5+, [FeFe]5+ or Fe5+, whose crystal structure is shown in Figure 1. In the structure, the metal centers are situated at the vertex of a triangular bipyramid where the bpp− acts as a bridging ligand between the axial and equatorial Fe, whereas the oxido group bridges the three Fe sites situated in the equatorial plane. The axial Fe sites are hexacoordinated (CN6) with a distorted Oh symmetry, whereas the equatorial ones are pentacoordinated (CN5) with a distorted C2v symmetry. This crystal structure is very similar to the previously reported one for Fe3+ complex (Okamura et al., 2016), except for a slightly shorter Fe-O and Fe-N distances as expected (Figure S1 and Table S1) (Sreerama and Pal, 2004).

It is important to realize here that the quantitative yield of this reaction implies that the pentanuclear structure is especially stable from a thermodynamic perspective given the large number of complexes that can be potentially formed by mixing the bpp− ligand and the iron [FeII(H2O)6]2+ (obtained from the dissolution of iron sulfate in water), as shown in Figure S2. These potential complexes range from simple mononuclear complexes with different number of bbp− ligands bonded to the Fe center to dinuclear, polynuclear, etc., again with different numbers of coordinated bpp− ligands attached to Fe. The formation of Fe3+ as the only complex generated in this reaction points to a scenario whereby the Fe-N bond is forming and breaking easily in agreement with the lability of high spin Fe(II) and Fe(III) complexes (Helm and Merbach, 2006). This is exemplified in Figure S2 where the first Fe complexes that will most likely be made upon mixing Fe(II) and bpp− are shown. The fac-[Fe(L)3]- complex will end up acting as a capping group for the final Fe3+ complex. Only the [Fe(L) (H2O)4]+ and the cis,cis-[Fe(L)2(H2O)2] complexes will have the proper ligand geometrical coordination so that they can lead to the formation of fac-[Fe(L)3]- without the need of additional ligand rearrangements. All the other complexes will need to rearrange, and thus Fe-N bonds will have to be broken and made anew to be able to lead to fac-[Fe(HL)3]-. Thus, all the potential complexes generated at the initial stages of the synthesis will reorganize in order to converge to the most thermodynamically stable complex, which is, in this case, Fe3+. A parallel phenomenon also occurs with other metal complexes such as Mn that can achieve a similar structural arrangement with bpp−, such as [MnII4MnIII(μ-bpp)6(μ3-O)]3+, Mn3+, and also with related linearly arranged tetranucleating ligands (Bao et al., 2010; Romain et al., 2011).

Although the Fe3+ is relatively stable in solution at low concentrations of water, the corresponding Mn analog, the Mn3+, decomposes almost immediately to generate the free ligand and [Mn(H2O)6]2+/3+ (Romain et al., 2011). This indicates the capacity of H2O to compete for the first coordination sphere of the Mn center, so that once a water molecule coordinates to a Mn center the whole structure collapses losing the stability provided by the pentanuclear arrangement.

In MeCN as solvent, the Fen+ complex is a very rich molecule from a redox perspective accessing six different oxidation states ranging from [FeII5(μ3-O) (μ-bpp)6]2+, Fe2+, where all the iron centers have oxidation state II up to [FeIII5(μ3-O) (μ-bpp)6]7+, Fe7+, where now all Fe centers have oxidation state III. All the oxidation states can be accessed by successive one electron electrochemically quasireversible processes, as can be observed in the CV in Figure 2 and in agreement with a previous report (Gouré et al., 2016) (see Figure S4 and Table S2 for further details). All CVs in this work are carried out using a glassy carbon (GC) electrode as a working electrode, an Ag+/Ag (0.01 M) as reference electrode and a Pt disk as auxiliary electrode unless explicitly mentioned. All potentials in this work are reported versus Fc/Fc+. The fully reduced species Fe2+ is air sensitive and thus needs to be isolated in an inert atmosphere. The fully oxidized species, Fe7+, displays a chemically reversible behavior during the CV timescale, but on bulk electrolysis timescales it is not stable and decomposes, indicating the high reactivity of such a high oxidation state species. All other Fen+ species in intermediate oxidation states are stable and can be isolated as solids in an open air atmosphere that is in accordance with a previous report (Gouré et al., 2016).

Figure 2

Electrochemical Characterization in MeCN of Fe3+

Cyclic voltammetry experiments for Fe3 + 0.2 mM dissolved in a 0.1 M TEAP MeCN solution (Vi = Vf = - 0.24 V; VC1 = −1.08 V; VC2 = 1.42 V) black trace, and in a 10:1 MeCN:H2O volume ratio (Vi = Vf = −0.24 V, VC1 = −1.08 V; VC2 = 1.19 V) red trace (background subtracted) at a scan rate of 10 mV/s. Labels indicate the oxidation state zones of predominance as a function of potential. The inset shows an enlargement of the 0.6–1.4 V zone.

See also Tables S1 and S2 and Figure S4.

The electronic structure of the six Fen+ (n = 2–7) complexes at the different oxidation states has been unambiguously established based on EPR and magnetic measurements (Gouré et al., 2016). The six coordinated apical Fe(II) centers in Fe have a low spin (LS) d6 configuration, whereas two equatorial Fe(II) have a high spin (HS) d6 and the third equatorial Fe(III) is a low spin d5. On the other hand, for the highest oxidation state complex, Fe7+, the apical Fe centers are LS, whereas the equatorial ones are HS.

Redox Properties in Aqueous MeCN

The Impact of [H2O] in the Catalytic Activity

In water, the Fe3+ complex is not soluble, but it can be solubilized in mixtures of MeCN and H2O. The latter is important since in the absence of water, the potential active species needed to enter into the water oxidation catalytic cycle cannot be formed. The electrochemical work reported here is carried out in mixtures of a MeCN solution containing 0.1 M tetraethyl ammonium perchlorate (TEAP) and water in a maximum 10:1 MeCN:H2O volume ratio. This from now on will be referred in an abbreviated manner as 10:1 MeCN:H2O.

In 10:1 MeCN:H2O as solvent mixture the CV of Fen+ is similar to the one reported in 0.1 M TEAP MeCN with a slight cathodic shift of roughly 40 mV for the wave associated with the [FeFe]6+/[FeFe]5+ couple as can be observed in Figure 2. Furthermore, a large electrocatalytic current starting at 0.85 V is observed that is assigned to the oxidation of water to dioxygen. This electrocatalytic wave was wrongly assigned to the catalytic activity of a molecular Fen+ complex (Okamura et al., 2016). It was proposed based on DFT that, once the [Fe]7+ species is reached, a solvent water molecule could coordinate in one of the equatorial iron centers forming [FeIII5(H2O) (μ3-O) (μ-bpp)6]7+, with increasing coordination number from CN5 to CN6. A series of oxidations and an additional water coordination to a neighboring Fe center was then proposed to occur so that a sufficiently reactive species would form an O-O bond that might finally release dioxygen. The transition state energy for the initial step, the coordination of water and the formation [FeIII5(H2O) (μ3-O) (μ-bpp)6]7+, is highly endergonic by 18.9 kcal/mol and is proposed to be the rate-determining step (rds) of this catalytic cycle (Liao et al., 2018). This is in agreement with the high stability of the Fen+ structure and thus the large energy needed to disrupt it. On the other hand, it also points out that once a water molecule coordinates a metal center, the whole structure might collapse with the formation of multiple Fe complexes containing different ratios of aqua and bpp− bonded ligands. If this disruption occurs in close proximity with a glassy carbon electrode the new species will generate FeOx as will be shown below.

Figure 3 left shows the third CV cycle of Fe3+ within the potential range of −1.08 to 1.19 V at a scan rate of 10 mV/s in 10:1 MeCN:H2O solution (the first cycle is reported in Figure S5A). The waves associated with the Fen+ molecular complex remain the same as in the first cycle (see Figure 2), whereas the catalytic current in the third cycle increases nearly five times from approximately 200 μA/cm2 up to approximately 1.0 mA/cm2 at 1.19 V. This unambiguously indicates the formation of new catalytically active species potentially adsorbed at the surface of the electrode. Indeed, transferring the glassy carbon electrode obtained after the CV into a clean electrolyte solution and scanning from 0.92 to 1.19 V shows a current density at 1.19 V that is close to 90% of the previous one. Furthermore, no redox waves associated with molecular Fen+ complex can be observed when scanning within the −1.08 to 1.19 V potential range (see Figure 3 left). These two experiments point out that FeOx are the main species responsible for the electrocatalytic activity observed here, given its well-known catalytic behavior (Le Formal et al., 2015). Furthermore, X-ray absorption spectroscopy (XAS) was also carried out on glassy carbon plates, which gives additional support to the formation of FeOx adsorbed at the electrode surface as will be discussed later on.

Figure 3

Electrochemical Analysis at Different Amounts of Water

Left, black trace, CV of the third cycle for Fe3 + 0.2 mM dissolved in a 10:1 MeCN:H2O solution (Vi = Vf = −1.08 V; VC1 = 1.19 V) at a scan rate of 10 mV/s. Red trace, CV of the GC working electrode obtained in the previous experiment immersed in a clean electrolyte solution. Gray trace, CV of a bare GC electrode under the same conditions. Inset, enlargement in the zone of the non-catalytic waves of the complex. Right, plot of current density at 1.19 V under different MeCN:H2O ratios obtained in the CV of: (1) black trace, after 2.5 cycles for a 0.2 mM Fe3+ solution with a GC electrode as WE at the previous conditions; (2) red trace, after the previous 2.5 cycles the GC electrode obtained is immersed in a clean electrolyte solution (Vi = Vf = 0.92 V; VC1 = 1.19 V); (3) gray trace, blank for a bare GC electrode. See also Figures S5A and S5B.

A series of related experiments were also carried out by changing the relative concentration of H2O from 1% to 10% in MeCN and are reported in the Figure S5B. In Figure 3 right a plot of the current density at 1.19 V versus the concentration of water is displayed for the initial Fe3+ solution and for the electrode obtained from this solution placed subsequently in a clean electrolyte solution. The very close values obtained here further point out that the Fen+ species are a precursor for the generation of FeOx that is actually the active catalyst. The difference between the initial current density and the one obtained in a clean electrolyte solution can be due to ligand oxidation, the formation of transient active species generated during the decomposition process to FeOx, or from the partial solubilization of the FeOx from the electrode.

As can be seen in Figure 3 right, the intensity of the wave at 1.19 V increases with the concentration of water thus clearly establishing a direct correlation between the H2O concentration and the catalytic activity. This points out to the presence of a series of equilibria between the initial Fen+ complex and FeOx as depicted in Figure 4. The larger the concentration of water, the larger is the equilibrium shift toward the aquated species, and thus a larger amount of FeOx will be deposited at the surface of the glassy carbon electrode. Furthermore, the increase in water concentration implies an increase in substrate concentration that can lead to a higher catalytic current.

Figure 4

Scheme for the Formation of FeOx from Fe53+

Potential non-isolated intermediate decomposition species formed from the Fen+ complex toward the generation of free iron, [Fe(H2O)6]n+, and the subsequent formation of the catalytically active species FeOx detected at the surface of an electrode.

The Influence of pH on the Stability of the Fe5n+ Complex

The stability of Fe complexes is strongly dependent on pH as has been shown in a number of occasions for related ligands (Draksharapu et al., 2012; Hong et al., 2013). The main driving force for decomposition process is the lability of the bonded ligands that can be substituted by solvent water ligands. In acidic pH this substitution process will be further enhanced by the protonation of the bonded ligands that will be strongly dependent on pH. In this respect, the pKa of pyridinium ion is 5.5 and that of pyrazole is 14.2. However, when the Hbpp is coordinated to a transition metal as in the case of [Ru(trpy)(Hbpp)]2+ (where trpy is 2,2’:6′,2″-terpyridine), then the pKa of the pyrazole moiety becomes more acidic with pKa values in the range of 5–7 depending on the oxidation state of the metal (Sens et al., 2003). In basic pH, the anionic OH− ligand will be responsible for the substitution process and subsequent generation of FeOx (Chen et al., 2013; Hong et al., 2013).

For water oxidation catalysis, the fact that every oxygen molecule evolved generates four protons implies that the local pH will also be strongly reduced in the double layer during water oxidation electrocatalysis. This will in turn strongly effect the stability of the complex leading to decomposition reactions at low local pHs. For this reason, it is extremely important to carry out water oxidation catalysis in the presence of a buffer so that the pH can remain practically constant. This strong influence of pH into the electrocatalytic activity is clearly manifested by observing the current density in the CV at 1.19 V for Fe3+ in the presence and absence of buffer (see Figure S6 right). Indeed, in the absence of a buffer, the current density observed is 44% larger than the one in the presence of a non-coordinating borate buffer that clearly suggests the influence of the local pH on the equilibria proposed in Figure 4.

CV experiments were also carried out at different pH values, and it was found that from pH = 2–7, the behavior of the Fen+ complex is basically the same (Okamura et al., 2016). However, below pH 2 the Fe3+ complex is not stable and decomposes to [FeII(H2O)6]2+ and free ligand as is the case of related complexes reported in the literature (Draksharapu et al., 2012; Hong et al., 2013). Furthermore, no electrocatalytic activity is observed at pH values below 2. (see Figure S7 for further details).

The Formation of FeOx Films at the Surface of the GC Electrode

The nature of the FeOx deposited at the surface of the GC electrode was evaluated by means of electrochemical, spectroscopic, and microscopy techniques. Figure 5 left shows the results of 100 repetitive CVs scans from −0.44 to 1.19 V for a 0.2 mM solution of Fe3+ in a 10:1 MeCN:H2O solution using a GC disk as a working electrode (GC). As it can be observed after the 100th cycle, the CV becomes nearly featureless with an increased double layer capacitance indicating that the initial electrode has lost its conductivity. A simply eye inspection of the electrode shows the formation of a film at the surface. Placing this electrode into a clean solution containing a ferrocene solution 0.2 mM shows that the anodic III/II wave has lost 93.3% (see Figure S8) of its area with regard to a pristine electrode in exactly the same conditions, confirming the isolating nature of the oxide deposited at the electrode.

Figure 5

Stability Analysis and Formation of FeOx Film

Left, 100 repetitive CVs for Fe3+ 0.2 mM dissolved in 10:1 MeCN:H2O (Vi = Vf = −0.24 V, VC1 = −0.44 V; VC2 = 1.19 V). Color code: black trace, first cycle; red trace, last cycle; gray traces, intermediate cycles displayed every five cycles. Right, blue trace, plot of the current density at 1.19 V as a function of CV cycles. Green trace, plot of ip,a of the [FeFe]6+/[FeFe]5+ redox wave as a function of CV cycles. See also Figures S8, S9, and S11.

The SEM image of this electrode does not show any boundary or particle shapes, but rather a homogeneous surface with similar morphology as the GC electrode, and thus we attribute this to the formation of a film. The nature of this film was further evaluated based on energy-dispersive X-ray (EDX) spectroscopy, scanning electron microscopy (SEM), and X-ray photoelectron spectroscopy (XPS) displayed in the Supplemental Information (Figures S9 and S11 top, respectively). XANES (X-ray absorption near edge structure) and EXAFS (extended X-ray absorption fine structure analysis) were also carried out on the Fen+ complex before (Figures S19-S21, Table S3) and after bulk catalysis (Figure 6) in a glassy carbon plate following exactly the same protocol used for the CV with the GC disk electrode. The XANES and EXAFS spectra shown in Figure 6 unequivocally show the spectral features of FeOx (Kuzmin and Chaboy, 2014; Tangwatanakul et al., 2017) at the electrode after the 100th cycle, thus discarding the potential surface absorption of the molecular Fen+ species. As observed by the red arrows in Figure 6B, a prominent increase in the amplitudes of the EXAFS peaks at apparent distances ∼1.5 and 2.5 Å are indeed observed in agreement with the EXAFS spectral features of Fe2O3 (shown in cyan).

Figure 6

X-ray Absorption Spectroscopy of the FeOx Electroactive Species

(A) Normalized Fe K-edge XANES of: (1) the Fe3+ complex in a frozen mixture of 10:1 MeCN:H2O (blue); (2) the frozen solution obtained after bulk electrolysis of Fe3+ at an applied potential of 1.42 V for 30 min (green; see main text and Supplemental Information for further experimental details); (3) glassy carbon plate (used as a working electrode) obtained after 100 successive CV experiments carried out in the ranges of −0.44 to 1.19 V of 0.2 mM solution of the Fe3+ complex in a mixture of 10:1 MeCN:H2O (brown); (4) Fe2O3 powder (cyan). (B) Experimental Fourier transforms of k-weighted Fe EXAFS of the samples described in the left using the same color code. The red arrows indicate the main peaks associated with Fe2O3. See also Table S3 and Figures S19–S21.

A closer inspection at the repetitive CV experiment provides additional insight into the progressive formation of the FeOx film. Figure 5 right shows a plot of the current density obtained at 1.19 V versus ip,a of the [FeFe]6+/[FeFe]5+ redox wave preceding the electrocatalytic current. During the first ten cycles, the electrocatalytic current at 1.19 V increases its intensity, whereas the ip,a of the [FeFe]6+/[FeFe]5+ wave decreases. This implies that the catalytic activity increases initially as active FeOx is being formed and deposited at the GC electrode. At the same time and owing to the formation of non-conductive FeOx (potentially due to a thicker layer or different morphology), the available conductive surface area decreases and thus the intensity of the molecular Fen+ waves decreases as well. The next 20 cycles show a decrease of both currents owing to the non-conductive nature of the film generated at the electrode. From 50 to 100 cycles, small amounts of isolating FeOx are further deposited, which generates a large increase of the capacitance at the electrode.

The Formation of FeOx Nanoparticles at the Surface of the GC Electrode

The performance of the Fe3+ complex was evaluated by multiple consecutive cycle voltammetric experiments under a different range of potentials. Figure 7 left shows 50 repetitive cyclic voltammetry experiments carried out in exactly the same conditions as the previous one but after changing the potential range from −1.08 to 1.19 V to fully reduce the initial complex all the way to the Fe2+ species. Figure 7 right also shows the current density plot obtained at 1.19 V versus ip,a of the [FeFe]6+/[FeFe]5+ redox wave. As the number of cycles proceed, the catalytic intensity at 1.19 V increases owing to the increasing amount of FeOx adsorbed at the GC electrode until it reaches a plateau due to the saturation of the surface. On the other hand, the intensity of the molecular Fen+ species waves decrease owing to a decrease of the concentration of the double layer caused by the formation of FeOx. As can be seen in the Figure 7, the FeOx deposited at the electrode is conductive. The formation of nanoparticles can be observed with SEM (see Figure S10) with an approximate diameter of about 40 nm.

Figure 7

Stability Analysis and Formation of FeOx Active Nanoparticles

Left, 50 repetitive CVs for Fe3+ 0.2 mM dissolved in 10:1 MeCN:H2O (Vi = Vf = −0.24 V, VC1 = −1.08 V; VC2 = 1.19 V). Color code: black trace, first scan; red trace, last scan; gray traces, intermediate cycles displayed every five cycles. Right, blue trace, plot of the current density at 1.19 V as a function of CV cycles. Green trace, plot of ip,a of the [FeFe]6+/[FeFe]5+ redox wave as a function of CV cycles. See also Figures S10 and S11.

The conductivity of the material also enables us to observe the formation of other electroactive species growing at the double layer (see cathodic waves at −0.25 and 0.75 V in Figure 7 left). This suggests that the initial Fen+ complex decomposes to other molecular species as indicated in Figure 4, before forming FeOx, which finally ends up being adsorbed at the surface of the electrode.

Potentiostatic Generation of FeOx

Potentiostatic experiments were carried out using a glassy carbon rod (GCrod) or indium tin oxide (ITO) as a working electrode as shown in the Supplemental Information (see Figures S12, S14, and S17).

A controled potential electrolysis (CPE) was performed with a GCrod as working electrode with 6.5 mL of a 0.2 mM (1.3 μmol) solution of Fe3+ and was carried out for 1 h at Eapp = 1.19 V. During this time 1.05 C was passed together with the formation of 2.5 μmol of O2 that accounts for 90 of faradaic efficiency (FE) (see Figures S12 and S13). Oxygen detection obtained through Clark electrode.

Potentiostatic experiments using ITO electrodes (S = 2 cm2) as working electrodes are shown in the Supplemental Information (Figures S14 and S17). Same conditions previously described, 6.5 mL of a 0.2 mM (1.3 μmol) solution of Fe3+ was applied. One CPE was carried out for 1 h at Eapp = 1.42 V. During this time 7.5 C (77 μmol of electrons/4 = 19.4 μmol of O2) was obtained, which corresponds to a TN = 14.9 assuming a 100% FE (calculated for comparison purposes; see Figure S14). After the bulk electrolysis the ITO electrode was placed in a clean electrolyte solution and it showed the same activity as in the presence of the Fe3+ solution, demonstrating again that the water oxidation activity is due to the formation of FeOx at the surface of the electrode (Figure S15 left). Furthermore, a CV using a GC disk as working electrode was carried out for the Fe3+ solution obtained after the bulk electrolysis and showed no molecular species present indicating that the whole solution is transformed to FeOx (see Figures S15 and S16). An additional bulk electrolysis was carried out at Eapp = 1.19 V (Figure S17) under the same conditions as in the previous case using an ITO electrode (S = 2 cm2), yielded 1 C (10.2 μmol of electrons/4 = 2.5 μmol of O2), which corresponds to a TN = 1.9 assuming a 100% faradaic efficiency. The CV of the solution after the CPE shows that a significant amount of the initial complex together with other waves associated with potential decompositions of the initial Fe3+ complex is still present (Figure S18 right). This implies that the initial complex has only been partially decomposed to FeOx and that this process is taking place slowly and in parallel to the electrocatalytic formation of O2 by the adsorbed FeOx.

Discussion

The compact structure of the Fe3+ complex is a highly stable structural arrangement and thus constitutes a thermodynamic sink in the sense that all intermediate species generated from the reaction of bpp− and Fe(II) can break and form new Fe-N bonds until they end up trapped in Fe3+. This implies that bond formation and breaking acts as an automatic healing process that leads to the final Fe3+ complex. A similar type of phenomenon has been described in supramolecular chemistry for the generation of macrocycles and cages based on other transition metals (Cook et al., 2013). This large stability is also displayed by related tetradentate linear ligands with Fe and other first row transition metal complexes such as Mn that generate virtually identical structures (Kabata-pendias and Mukherjee, 2007). The large degree of stability is also manifested in the large energy value (18.9 kcal/mol) calculated that is required to coordinate an additional water molecule to one of the equatorial Fe center, changing its first coordination sphere from CN5 to CN6.

These pentanuclear complexes are very stable from a thermodynamic perspective, but they also are highly labile owing to their high spin electronic configurations. Thus, although these complexes are stable in MeCN solution, they readily decompose in the presence of coordinating solvents such as water, where the latter competes for the first coordination sphere. This decomposition phenomenon involves the disruption of the whole complex leading to a structure crumbling effect that finally generates the [M(H2O)6]2+/3+ and the free ligand as has been reported for the Mn3+ case. The Fe3+ complex is slightly more stable than its Mn analog and can tolerate concentrations up to 10:1 MeCN:H2O at low oxidation states. However, the stability of the complex is reduced in its high oxidation state species even in MeCN. Although the [FeFe]6+ decomposes completely in about 1 h, the [Fe]7+ decomposes much faster and has not been isolated (Gouré et al., 2016). Electrochemically, the [FeFe]6+ species are not stable in a 10:1 MeCN:H2O solution as ascertained by CV leading to the aquated species (Figure 4). Additionally, increasing the applied potential to the zone of predominance of the [Fe]7+ species leads to the formation of FeOx concomitant with the generation of a large electrocatalytic water oxidation current. The fact that the foot of the electrocatalytic current is found at a 150 mV lower potential than the foot of the [Fe]7+/[FeFe]6+ couple suggests that both the [FeFe]6+ and the [Fe]7+ might be responsible for the aquation of the Fen+ species that leads to the formation of FeOx adsorbed at the electrode. Given the large stability of the Fen+ structure, once a water adds to the first coordination sphere of an equatorial Fe center, the whole structure immediately crumbles giving rise to a large number of potential species as outlined in Figure 4. This view is also in agreement with the increased catalytic activity obtained upon increasing the H2O concentration, which shifts the equilibria to the right as shown in Figure 3. These molecular high-oxidation-state species generated at high potentials from the dismantling of the Fen+ structure could potentially behave as water oxidation catalysts. However, in the present case, given the large activity associated with FeOx, the activity of the resulting decomposed species is small if not active at all, as evidenced by CV experiments in clean electrolyte solutions.

All these experiments suggest the presence of a very delicate equilibrium between the [FeFe]6+ and the [Fe]7+ species, and their aquated counterparts as proposed in Figure 4. An additional evidence of this delicate equilibrium is exemplified by the experiments carried out in the presence of triflic acid that completely suppresses catalysis. Furthermore, the experiments carried out at different pHs suggest the local pH close to the electrode can reach pH values as low as 1, in experiments carried out in the absence of a buffer, for instance, in a 10:1 MeCN:H2O solvent. The low pH conditions could further help in the aquation of the Fe3+ complex and thus in the generation of FeOx.

An additional interesting point that also emerges from this work is how the nature of the FeOx formed at the surfaces of the electrode (conductive versus isolating; films versus nanoparticles), is strongly dependent on the potential range used to generate it. Furthermore, it is also striking to see the high activity of this FeOx adsorbed at the surface of the electrode that reaches current densities in the range of 3 mA/cm2 at pH 7.

In this regard, the high activity of the FeOx could lead to a misinterpretation of the results if the working electrodes are not properly evaluated in clean electrolyte solutions, since only very small amounts of the initial molecular complex are needed to be transformed into highly active FeOx. Thus, in the hypothetical case that the stability of the complex in solutions after a bulk electrolysis experiment was checked, for instance, by UV-vis, MS, or DLS spectroscopy, it would appear as if the initial catalyst was intact as the initial concentration would remain practically the same.

Several main conclusions can be inferred from the present work. In the first place, the auxiliary ligands used in WOCs with transition metals should contain oxidatively robust ligands given the high redox potentials of this reaction. Therefore, ligands containing benzyl pyridyl groups will be easily oxidized as has been shown in many cases (Radaram et al., 2011; Sander et al., 2015; Wang et al., 2016). Thus, they should not be used in their ligand framework. Second, special attention should be given to the stability of first row transition metals in water given the high lability of the M-L bond in this solvent (Helm and Merbach, 2006). The aqua substitution will foster the formation of oxides adsorbed at the surface of the electrode. Finally, the fact that the water oxidation reaction generates four protons per oxygen molecule implies that a buffer should always be used to avoid ligand decoordination and oxide formation.

Limitation of the Study

No limitation of the study can be declared.

Resource Availability

Lead Contact

Further information and requests should be directed to the Lead Contact, Antoni Llobet (allobet@iciq.cat).

Materials Availability

No new reagents were synthetized. There are no restrictions to the availability of chemicals.

Data and Code Availability

Crystallographic information for [FeFe]5+ with CCDC number 1963878 is available at https://www.ccdc.cam.ac.uk/.

Methods

All methods can be found in the accompanying Transparent Methods supplemental file.