An Iron Electrocatalyst for Selective Reduction of CO2 to Formate in

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An Iron Electrocatalyst for Selective Reduction of CO2 to Formate in Water: Including Thermochemical Insights Atefeh Taheri, Emily J. Thompson, James C. Fettinger, and Louise A. Berben* Department of Chemistry, University of California, Davis, California 95616, United States S Supporting Information *

ABSTRACT: C−H bond formation with CO2 to selectively form products such as formate (HCOO−) is an important step in harnessing CO2 emissions as a carbon-neutral or carbonnegative renewable energy source. In this report, we show that the iron carbonyl cluster, [Fe4N(CO)12]−, is an electrocatalyst for the selective reduction of CO2 to formate in water (pH 5− 13). With low applied overpotential (230−440 mV), formate is produced with a high current density of 4 mA cm−2 and 96% Faradaic efficiency. These metrics, combined with the long lifetime of the catalyst (>24 h), and the use of the Earthabundant material iron, are advances in catalyst performance relative to previously reported homogeneous and heterogeneous formate-producing electrocatalysts. We further characterized the mechanism of catalysis by [Fe4N(CO)12]− using cyclic voltammetry, and structurally characterized a key reaction intermediate, the reduced hydride [HFe4N(CO)12]−. In addition, thermochemical measurements performed using infrared spectroelectrochemistry provided measures of the hydride donor ability (hydricity) for [HFe4N(CO)12]− in both MeCN and aqueous buffered solution. These are 49 and 15 kcal mol−1, respectively, and show that the driving force for C−H bond formation with CO2 by [HFe4N(CO)12]− is very different in the two solvents: +5 kcal mol−1 in MeCN (unfavorable) and −8.5 kcal mol−1 in water (favorable). KEYWORDS: CO2 reduction, iron, electrocatalysis, carbonyl cluster, hydricity, formate

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(Table 1).10 The Ir catalyst was also mobilized on carbon nanotubes (CNTs) and showed similar activity for formate production.11 Recently, Kanan and co-workers improved the performance of Pb electrodes using nanocrytalline Pb prepared

INTRODUCTION Historical records show correlations between average atmospheric temperature and CO2 concentration, and this is one reason that CO2 is implicated as a contributor to currently observed patterns of global temperature increase.1 The efficient reduction of CO2 to a C−H bond-containing product for use as a chemical fuel has the potential to establish a CO2 neutral economy where no additional CO2 is emitted by human endeavors.2,3 Catalytic systems are known that convert CO2 selectively to CO.4 However, a significant challenge facing catalyst development for CO2 reduction is selectivity in product formation, when C−H and C−C bond-forming reactions are involved. As an example, heterogeneous Cu metal electrodes produce C−H and C−C bonds impressively, but mixtures of H2, CO, formate, ethylene, and methane are obtained.5 In recent work by Kanan et al., a copper electrode annealed in pH 7.2 buffer solution produced formate and CO with Faradaic efficiency (FE) values of 33% and 40%, respectively. In addition, most molecular electrocatalysts that produce formate also produce CO or H2, or both, and often also oxalate.6 The importance of generating higher selectivity in these reactions to produce C−H and C−C bond-containing products from CO2 in aqueous solution is further underscored by the recent high levels of activity in the area.7−9 In one case, an Irbased molecular electrocatalyst was reported to operate at 0.6 mA cm−2 at −1.65 V vs SCE in organic and aqueous solutions © XXXX American Chemical Society

Table 1. Electrocatalytic Formate Production from CO2 in Aqueous Solutions

ref

catalysta

10b 11

IrHL(S)2 IrHL(S)2 carbon nanotubes Pb Pd on Pt 10% Pd on C SnO2 nanocrystals [Fe4N(CO)12]−

13 14 16 18 this work

applied potential (V vs SCE), E

Faradaic yield, formate [%]

current density, j [mA cm−2]

stabilityb [h]

−1.65 −1.65

93 83

0.6 3.5

25 2

−1.4 −0.83 −0.77 −1.8

98 100 99 93

0.3 0.3 4 5.5

95% FE for formate production with a maximum current density of 0.3 mA cm−2 at −0.83 V vs SCE.14 Koper and co-workers also reported formate production using a Pd−Pt electrode at low overpotential; however, efficiency was not reported.15 Kanan and co-workers greatly improved the performance of that material by loading Pd onto C so that formate was produced selectively at just −0.77 V vs SCE.16 Lastly, several reports have demonstrated formate production by reduction of CO2 at tin-based electrodes.17 In particular, nanoparticles of SnO2 have been shown by Meyer and co-workers to enable formate generation at −1.8 V vs SCE from aqueous solutions.18 Of particular concern to us is the development of highly stable Earth-abundant catalysts for formate production, and also a detailed mechanistic and thermochemical understanding of catalyst properties to improve catalyst design. Our efforts to achieve these goals have been summarized so far in two reports. Initially, we showed that an iron carbonyl cluster is an electrocatalyst for hydrogen evolution or formate formation, depending on the reaction conditions in MeCN solutions.19 Only trace formate was detected in that study. We subsequently reported a series of iron carbonyl clusters and each were electrocatalysts for H2 evolution from water at pH 5.20 We identified [Fe4N(CO)12]− as the most stable of these electrocatalysts, and at pH 5, it afforded H2 with 95% FE at −1.25 V vs SCE. In this report, we show that [Fe4N(CO)12]− (1−) is a selective and stable catalyst for C−H bond formation with CO2 to yield formate in aqueous solution (Chart 1). Between pH 5

Figure 1. (Left) Cyclic voltammograms of 1 mM 1− in 0.1 M Bu4NPF6 MeCN solution (black), and in MeCN/H2O (95:5) (blue). Scan rate: 100 mV s−1. (Right) Cyclic voltammograms of 0.25 mM 1− in 0.1 M Bu4NPF6 MeCN/H2O (95:5), at different scan rates, under N2. Inset: jp vs υ0.5 for reduction. Overlaid line is the linear fit. Glassy carbon working electrode.

catalyst was varied over a series of CV experiments with 1−, we observed that the addition of waterup to 5% of the total solution volumeresulted in very minimal changes in current or peak potential for the two reversible events at −1.23 and −1.6 V (Figure 1, left). A new oxidation event at −0.5 V did appear, which we have previously proposed, corresponds to oxidation of the hydride intermediate (H-1)−.19 This oxidation event is present even when the scan direction is reversed as early as −1.1 V (see Figure S1 in the Supporting Information). To further support our hypothesis that H2 evolution is not occurring appreciably when MeCN/H2O (95:5) solutions of 1− are electrolyzed at −1.2 V, we performed CV analysis with scan rates varied from 2 mV s−1 up to 1 V s−1 (Figure 1, right). The plot of cathodic current density versus υ1/2 is a straight line consistent with a diffusion-controlled one-electron event, as described by the Randles−Sevcik equation (eq 1):22

Chart 1. Line Drawing for [Fe4N(CO)12]−

ip = (2.686 × 105)n3/2AD1/2Cυ1/2

(1)

where n is the number of electrons (here, n = 1), A the electrode area (A = 0.0707 cm2), D the diffusion coefficient for the complex (expressed in units of cm2 s−1), C the concentration of the complex (C = 0.2 mM), and υ the scan rate (expressed in units of V s−1). The straight line is not consistent with a catalytic event on the CV time scale (Figure 1, right). In the absence of CO2, H2 is evolved on the longer time scale of a controlled potential electrolysis (CPE) experiments.19 From the slopes of the plot of jp vs υ1/2, the diffusion coefficient of 1− was calculated to be ∼0.8 × 10−5 cm2 s−1. Cyclic Voltammetry Analysis of 1 − in Aqueous Solution under N2. We have previously reported the pHdependent CV responses for 1− between pH 5 and 9,20 where 1− has a pH-independent catalytic reduction event at −1.25 V. In CPE experiments, 1− is a catalyst for reduction of buffered water to H2 with near-quantitative FE and modest operating potential (−1.25 V). Electrocatalytic CO2 Reduction by 1− Selectively to Formate in MeCN/H2O (95:5). When CO2 was added to CV experiments performed with 1− in MeCN/H2O (95:5), a slight increase in current was detected at −1.2 V, along with a loss in reversibility of the peak (Figure 2, left). CPEs of 0.1 mM 1− solutions saturated with CO2 were performed at −1.2 V vs SCE (Figure S2 in the Supporting Information) and both the headspace and solution were analyzed for products. H2 was detected in the headspace with 24 h. In the presence of CO2, water does not protonate (H-1)− appreciably, even at pH 5. We rationalize these results using thermochemical measurements conducted using infrared spectroelectrochemistry (IR-SEC) to characterize the equilibria associated with protonation of the catalyst.

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RESULTS AND DISCUSSION CV’s of 1− in MeCN/H2O (95:5) under N2. Electrochemical measurements for 1− were initially performed in the absence of CO2. Cyclic voltammetry (CV) experiments were performed on 1.0 mM solutions of 1− in dry 0.1 M Bu4NPF6 MeCN (Figure 1, left). As previously reported, CV analysis of 1− revealed two reversible couples, at E1/2 = −1.23 V and −1.6 V vs SCE, and these correspond to the 1−/2− and 12−/3− couples, respectively. Irreversible oxidation of 1− (not shown) occurs at potentials more positive than 0.2 V. In MeCN solution, the pKa of water is estimated to be in the range of 38− 41.21 When the concentration of water in MeCN solutions of 7141

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Table 2. CPE with 0.1 mM 1− in Buffered Aqueous Solutions (pH 5−13) Faradaic Yield

qa [C] pH pH pH pH pH pH pH

Figure 2. Cyclic voltammograms of (left) 1 mM 1− in 0.1 M Bu4NPF6 MeCN solution under N2 (black), under CO2 (red); and (right) 0.1 M KHCO3 with no catalyst at pH 6.5 (black), with 1 mM 1− in 0.1 under CO2 at pH 6.5 (red). Glassy carbon working electrode, 100 mV s−1.

5 6 6.5 7 8 9 13

21 16 19 45 37 19 10

± ± ± ± ± ± ±

1 1 1 1 2 1 1

HCO2− 80 91 96 95 86 91 95

± ± ± ± ± ± ±

3 0.5 2 5 1 1 1

turnover number, TONb (HCO2−)

H2 14 5 2 4 10 8 3

± ± ± ± ± ± ±

0.5 1 1 0.5 2 2 1

36 30 36 88 67 37 19

± ± ± ± ± ± ±

0.5 2 2 5 1 1 2

turnover frequency, TOF (HCO2−) [h−1] 43 36 44 106 80 44 23

± ± ± ± ± ± ±

0.5 2 3 6 1 2 2

Charge passed during 50 min of CPE. bTON is moles of HCO2− per mole of catalyst. No CO was detected. Applied potential was −1.2 V vs SCE. Each experiment repeated at least three times.

a

chromatography−thermal conductivity detection (GC-TCD), and formate was detected in the solution in 94% ± 3% FE, and quantified using 1H NMR spectroscopy. No CO was detected. The CPEs were operated at a current density of ∼0.7 mA cm−2 and a turnover frequency (TOF) of 15 ± 3 h−1. A control CPE experiment containing no 1− that was performed at −1.2 V also afforded small amounts of H2 at a very low rate of reaction (maximum 8 C passed over 50 min), so we ascribe the production of H2 in experiments containing 1− to background reduction of H+ at the glassy carbon working electrode. To provide further insight into the mechanism of CO2 reduction by 1− in MeCN/H2O (95:5), we performed CV experiments under 1 atm CO2 where the direction of scan was reversed at different potentials (see Figure S3 in the Supporting Information). When the scan was reversed between −1 V and −1.3 V, the event for oxidation of (H-1)− appeared at −0.4 V. The intensity increased with increasing potential for scan reversal until −1.3 V when the intensity decreased, because (H1)− reacts with CO2. In the absence of CO2, the same oxidation of (H-1)− was observed at −0.5 V (see Figure S1). These results are consistent with production of (H-1)− in the presence of CO2, and not with reaction of 1− or 12− with CO2 to afford a metallocarboxylate intermediate. Electrocatalytic CO2 Reduction by 1− Selectively to Formate in Aqueous Solution (pH 5−13). In aqueous buffered solutions saturated with CO2 (pH 5−13), an increase in current is observed relative to scans performed in the absence of 1− (see Figure 2 (left), and Figure S4 in the Supporting Information). The buffers employed were acetate (pH 5), KHCO3/HCO3− (pH 6.5), phosphate (pH 7), borate (pH 9), and sodium hydroxide−potassium chloride (pH 13). Under N2, 1− is an electrocatalyst for proton reduction, so comparison of the scans measured under CO2 and N2 does not provide a measure of catalytic current enhancement for CO2 reduction (Figure S4). CPE experiments were performed at −1.2 V on 0.1 mM solutions of 1−, and our analyses of the headspace and solution uniformly revealed selective formation of formate by 1− (see Table 2, as well as Figures S2, S5, and S6 in the Supporting Information). Small amounts of H2 were produced in a background reaction with the GC electrode. Using GC-TCD, no CO or CH4 were detected (Figures S7−S10 in the Supporting Information). Neither methanol nor oxalate was observed via 13C NMR spectroscopic analysis of CPE experiments performed with 13CO2 (Figure S11 in the Supporting Information). No methanol was detected by 1H

NMR spectroscopy (Figure S12 in the Supporting Information). Experiments performed in the absence of catalyst passed very little charge (Figure 3, left). When the GC working

Figure 3. (Left) Charge versus time plots for CPE experiments with 0.1 mM 1− (−1.2 V vs SCE) under CO2 in MeCN/H2O (95:5) and in buffered aqueous solutions (pH 5−13) over 50 min (lines with data points). CPEs recorded in the absence of catalyst, or performed with the uncleaned electrode in fresh electrolyte after a CPE with catalyst (solid lines). (Right, black) Current density versus time for CPE with 1− at pH 7 over 24 h (decrease in current at the end is due to the change in pH as buffer was depleted). (Right, blue) Faradaic efficiency (FE) of formate production versus time.

electrode used in a CPE experiment with 1− was reused in a solution containing only electrolyte (no catalyst), no formate generation was observed. This confirms catalyst plated onto the glassy carbon electrode does not produce formate (Figure 3, left). This is true even after CPEs were performed for over 24 h (as in Figure 3, right). The rate of reaction observed by CPE is pH dependent (Table 2), as observed from both the current density in CPE experiments and in the turnover number (TON) for formate production during the 50 min CPE experiments. This TON is an underestimation, because only catalyst molecules that are in the diffusion layer are reduced by the electrode and promote catalysis. Based on an analysis method proposed by Savéant and co-workers,23 we approximate that the amount of active catalyst in the CPE experiments is ∼3500 times smaller than the bulk solution, and thus the TON may be 3500 times larger than the numbers we report in Table 2. Details of this calculation are shown in the Supporting Information. The fastest rates of reaction were observed at pH 7, where the CPE experiments operated with an average current density 7142

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ACS Catalysis of 4.0 mA cm−2, with formate produced selectively with 95% ± 5% FE, for over 24 h (Figure 3, right). Infrared (IR) spectra before and after CPE also showed that the complex was stable for at least this 24 h (see Figure S13 in the Supporting Information). With the applied potential of −1.2 V vs SCE, the applied overpotential for formate production from CO2 falls from 440 mV at pH 7 down to just 230 mV at pH 13. This could explain the observed decrease in reaction rate. In aqueous solution, under 1 atm of CO2, we repeated the experiment where the direction of the CV scans was changed at different potentials (see Figure S14 in the Supporting Information). As in MeCN/H2O (95:5) solution (vide supra), we observed the formation of (H-1)− in the presence of CO2 when the scan direction was reversed between −0.8 V and −1.2 V. Scan reversal more negative than −1.2 V gave diminished quantities of (H-1)−, because (H-1)− reacts with CO2 to afford formate. As a further comment, in the absence of CO2 in aqueous solution (vide supra), no formate is detected as a product of electrolysis and this rules out a mechanism for formate formation where OH− anions generated by consumption of H+ react with the carbonyl ligands on the cluster. Mechanistic Investigations Using CV. The reduction of CO2 to formate catalyzed by 1− was further investigated using CV experiments in MeCN/H2O (95:5) and various buffered aqueous solutions. We initially assumed that the relationship in eq 2 is correct in order to determine the order of the reaction, with respect to [1−] and [CO2]:23,24 jcat,max = 2F[Cat] Dcat kcat[CO2 ]

Knowing that the reaction is first order in both catalyst and CO2, we next fit j versus potential to the relationship shown in eq 3 (Figure 4, left):23,24 jcat =

2F[Cat] Dcat kcat[CO2 ] F 0 ⎤ 1 + exp⎡⎣ RT (E − Ecat )⎦

(3)

where jcat is the current density of catalytic process (A cm−2), R the gas constant (R = 8.314 J K−1 mol−1), T the temperature (K), E the applied potential (V), E0cat the reduction potential of the catalyst (V). All of the other terms in eq 3 were previously defined for eq 2. For these analyses, [CO2] = 0.28 M in MeCN under 1 atm CO2 was used.25 The diffusion coefficient was acquired from the earlier scan rate study in MeCN/H2O (95:5) in the absence of CO2 (Figure 1). The current density at different concentrations of catalyst was fit to eq 3 (Figure 4 left, overlaid dotted lines), and these fits match the experimental data from the foot of the wave to the maximum current, which indicates that (1) the system is Nernstian, and (2) our original assumption that we can use eq 2 for data collected in MeCN/ H2O (95:5) is valid. Using the slope of the plot of jcat,max vs [1−] (Figure 4, left inset) and eq 2, the rate constant was determined to be 73 M−1 s−1 in MeCN/H2O (95:5).26 Using kcat = 37 M−1 s−1 and [CO2] = 0.28 M, we estimated that the turnover frequency (TOF) for the reduction of CO2 to formate in MeCN/H2O (95:5), is 10 s−1, as described by eq 4: TOF = kcat[CO2 ]

(2)

(4)

The rate of reaction for formate formation in aqueous solution could not be determined, because, in the absence of CO2, 1− is an efficient electrocatalyst for proton reduction; therefore, we could not measure a diffusion coefficient in aqueous solutions in the absence of CO2. Comparison of the CPE experiments recorded in MeCN/H2O (95:5) with CPE experiments in water (pH 7) gives a rough estimate that the rate of formate production is 7 times faster in aqueous solution. Based on the TOF value of 10 s−1 in MeCN/H2O (95:5), and the CO2 concentration of 33 mM in water under 1 atm CO2,27 we estimate a kcat value of 2121 M−1 s−1 in water. Characterization of 1− and Reduced Analogues Using Infrared Spectroelectrochemistry (IR-SEC) in MeCN. The one-electron reduced catalyst, 12−, is a proposed intermediate in the reduction of CO2 to formate, so we generated and characterized the IR spectra and the reactivity of these species using IR-SEC to learn more about the reaction mechanism. Initially, a solution of 1− in 0.1 M Bu4NPF6 MeCN was purged with nitrogen gas and injected into the thin-layer IR-SEC cell. At the resting state, 1− has two νCO stretches at 1989 and 2018 cm−1 (Figure 5, left).28 When a voltage is applied at the potential of the one-electron reduction seen in CV experiments, −1.4 V vs SCE, the growth of three new νCO stretches occurred at 1918, 1940, and 1962 cm−1, concomitant with the decay of the bands at 1989 and 2018 cm−1, with an isosbestic point at 1970 cm−1. These three new bands are assigned to the reduced catalyst [Fe4N(CO)12]2−, 12−. The observed shift of νCO to lower energy is consistent with the increase in electron density on the cluster which accumulates in the π* orbitals of the CO ligands.29 When a voltage was applied at a potential corresponding to the second one-electron reduction event observed by CV (−1.8 V vs SCE), 12− is reduced to 13− and a new broad band at even lower energy, 1879 cm−1, was observed (see Figure S17 in the Supporting Information).

where jcat,max is the maximum current achieved in CV experiments, F the Faraday constant (expressed in units of C mol−1), [Cat] the bulk concentration of catalyst (in units of mol cm−3), R the gas constant (R = 8.314 J K−1 mol−1), D the diffusion coefficient (in units of cm2 s−1), kcat the second-order rate constant of catalytic reaction (in units of mol−1 cm3 s−1), and [CO2] the concentration of CO2 in solution (in units of mol cm−3). In either MeCN/H2O (95:5) or buffered aqueous solutions saturated with CO2, the catalytic peak current density (jcat) varies linearly with [1−], which is consistent with a mechanism for CO2 reduction that is first-order in catalyst (Figure 4). The dependence of the current density on [CO2] indicated that the reaction is first order in [CO2] (see Figure S15 in the Supporting Information).

Figure 4. Cyclic voltammograms with varied [1−]. (Left) In 0.1 M Bu4NPF6 MeCN/H2O (95:5) under CO2. Overlaid dotted lines are the individual fits to eq 2. Inse shows jmax vs [1−]. (Right) In 0.1 M KHCO3 H2O under CO2 (pH 6.5). Inset shows jmax vs [1−]. Glassy carbon electrode, 100 mV s−1. 7143

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structure of the unreduced HFe4N(CO)12 (1) but did not locate the hydride, because of disorder.30 Instead, the hydride position was assigned based on an elongated bond length at the butterfly hinge. The structure of 12−/(H-1)− now shows that the reduced hydride is not simply a reduced analogue of the previously proposed structure of H-1, but that the hydride most likely bridges a wing edge of the butterfly structure. We were not able to resolve the bond lengths and angles for 12− and (H1)−; therefore, we do not discuss these in detail. We also probed the reaction of CO2 with 12−/(H-1)− and determined that reaction with CO2 produced formate in 25% yield, based on (H-1)− (see Figure S19 in the Supporting Information). This experiment further suggests that (H-1)− is responsible for C−H bond formation to afford formate. We could not quantify the possible formation of H2 (see the Experimental Section for more details). In Situ Characterization of CO2 Reduction to Formate by 1− in MeCN. As a further probe of the mechanism of CO2 reduction to formate by 1−, we followed the reaction using IRSEC in MeCN/H2O (95:5) at −1.4 V. Upon reductive electrolysis of the CO2-saturated solution, growth in the intensity of the broad band of νCO stretches ranging from 1970 cm−1 to 1800 cm−1 was observed concomitant with decay of bands at 1989 and 2018 cm−1. These observations indicate that 1− is disappearing while (H-1)− is formed (Figure 7, right). A

Figure 5. IR spectral changes during the reduction of (left) 1.2 mM 1− at −1.4 V vs SCE in 0.1 M Bu4NPF6 MeCN (inset shows difference absorption spectra upon reduction) and (right) FTIR spectra of 1− in THF (black trace), and after the addition of 1 equiv of Cp2*Co to give 12− (red trace).

Synthesis of Catalytic Intermediates. To further characterize the catalytic intermediates, we attempted to isolate 12− and (H-1)−. These compounds react quickly with water and proton sources, so these were avoided. The addition of 1 equiv of decamethylcobaltocene (Cp2*Co) to a THF solution of 1− afforded a species with the IR spectrum matching the difference spectrum we assigned as 12− from IR-SEC experiments (Figure 5, right). This product was crystallized as brown, block-shaped single crystals at −25 °C over a period of 4 months by diffusion of ether into a THF solution. The solid-state structure revealed that the crystals consisted of 50% 12− and 50% (H-1)−, with [Cp2*Co]+ as a countercation (see Figure 6, as well as Tables S1 and S2 in the Supporting

Figure 7. (Left) IR-SEC for the reduction of 1− at −1.4 V in 0.1 M Bu4NPF6 MeCN/H2O (95:5) performed under 1 atm CO2. Black arrows indicate a decrease or increase of bands. (Right) Normalized spectra from left overlaid with spectrum of (H-1)− in black. Band with red arrows are attributed to (HCOO-1)−.

Figure 6. Solid-state structure of 12−/(H-1)− in [Cp2*Co]3[HFe4N(CO)12][Fe4N(CO)12]. Green, gray, blue, red, and white ellipsoids represent Fe, C, N, O, and H atoms, respectively. Ellipsoids shown at 50%.

band for CO at 2138 cm−1 was not observed, which was consistent with our analysis of products from CPE experiments.31 Two new bands at 1684 and 1596 cm−1 were observed for formate and HCO3−.32 Control experiments performed with IR-SEC in the absence of catalyst showed no change in the IR spectrum over time. This indicates that, under our experimental conditions, at just −1.4 V vs SCE in MeCN, no formate formation occurred at the Au electrode. Au electrodes do catalyze CO2 reduction, with CO as the main product. However, this occurs in aqueous solutions at modest potentials, or from nonaqueous solutions at potentials greater than −2.3 V vs SCE.33,34 Bands for (H-1)− are very broad and could mask another species in the same region of the spectrum. Therefore, we looked for an intermediate species that could be assigned as a CO2 insertion product or outer-sphere interaction of CO2 with (H-1)−. The shoulder at 1958 cm−1 (Figure 7, right) is not present in the spectrum of (H-1)−, and, to aid assignment of this as a potential formate adduct (HCOO-1)−, we performed an experiment under 13CO2 (see Figure S20 in the Supporting

Information). Apparently, over the 4 month crystallization period, 12− reacted with trace water in the THF solution to afford 50% conversion to (H-1)−. An IR spectrum of the crystals revealed both 12− and (H-1)−, and no 1− (see Figure S18 in the Supporting Information). We have not yet been able to crystallize a pure sample of 12−. The butterfly geometry of 1− is maintained in the structure of 2− 1 /(H-1)−. The presence of a hydride in 50% of the structure was initially suspected to balance the charge: there are three [Cp*2 Co]+ cations present for every two clusters molecules in the unit cell. The presence of the hydride was confirmed by IR spectroscopy, and then located by analysis of the difference map, which indicated electron density along the butterfly wing between Fe2 and Fe3. The thermal parameter for the hydride was refined to U = 0.065 (see the Experimental Section for more details). In 1980, Muetterties and co-workers published a 7144

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ACS Catalysis Information): the feature at 1958 cm−1 disappeared, which is suggestive of a (HCOO-1)−-containing intermediate. However, we could not definitively locate the expected band for (H13COO-1)− at lower energy. The bands for formate and HCO3− also disappeared and new bands at 1632 and 1558 cm−1 appeared, consistent with the expected shifts for different carbon isotopes. We also attempted to perform IR-SEC experiments on aqueous buffered solutions. However, our IRSEC cell contains a Au working electrode and bubbles of H2 prevented the collection of useful IR-SEC spectra. Taken together, the experimental data in MeCN/H2O (95:5) and in aqueous solution point to a mechanism for CO2 reduction where the electrocatalyst is reduced from 1− to 12− and then protonated to afford (H-1)−. Analysis of the solution and headspace indicates that (H-1)− reacts selectively with CO2 to form a C−H bond to give formate (see Scheme 1).

[HFe4N(CO)12 ]− ⇌ [Fe4N(CO)12 ]2 − + H+ (H − 1)−

12 −

(5)

We first characterized a solution of (H-1)− using an IR-SEC experiment so that this spectrum could be used to quantify (H1)− in solutions containing the equilibrium mixture shown in eq 5. When 1− was electrolyzed in the IR-SEC cell at −1.4 V, in a solution of MeCN/H2O (95:5), quantitative formation of (H1)− was observed, as indicated by νCO stretches at 1937 and 1921 cm−1 (Figure 8, left). No 12− was detected.

Scheme 1. Proposed Mechanism for Reduction of CO2 to Formate by 1− in the Presence of Protons

Figure 8. (Left) IR-SEC spectrum at −1.4 V in 0.1 M Bu4NPF6 MeCN/H2O (95:5), showing the reduction of 0.6 mM 11− to (H-1)−. Inset shows normalized difference spectra. (Right) IR-SEC spectra in 0.1 M Bu4NPF6 MeCN at −1.4 V showing reduction of 0.65 mM 1− in the presence benzoic acid. Normalized difference spectra are presented in the inset, showing a decrease of the peak at 1940 cm−1, 12−, concomitant with increase of shoulder at 1920 cm−1, (H-1)−.

We know that (H-1)− is not protonated further by water, because no gas evolution (H2) was apparent. In addition, all normalized difference spectra were exactly the same, which confirms complete conversion of 1− to (H-1)− (Figure 8, left inset). IR stretches of metal hydrides and CO bond absorptions sometimes overlap,37 so the same IR-SEC experiment was performed in D2O. However, no change to the IR spectrum was observed (see Figure S21 in the Supporting Information), and this indicates that the observed bands likely correspond to νCO stretches and not to νFeH stretches.38 The thermochemical cycle shown in eqs 6−9 was used to determine the pKa of (H-1)− in MeCN. To measure Keq for eq 6, we reduced 1− to 12− via the addition of 1.1 equiv of electrons, in the presence of the weak acid, benzoic acid (pKa(PhCO2H) = 20.7 in MeCN),39 and monitored the formation of products using IR-SEC (Figure 8, right). A peak at 1940 cm−1, 12−, with a shoulder at 1920 cm−1, (H-1)−, appeared upon application of −1.2 V potential. The intensity of the shoulder, (H-1)−, increased over time. Subsequent scans show that the ratio of 12− to (H-1)− does not change over time, so the system is at equilibrium (Figure 8, right).

Thermochemical Measurements. The selective production of formate by the intermediate (H-1)− suggests that there is a high kinetic barrier toward reaction of that intermediate with H+. To further understand the thermochemistry of the catalytic cycle, we performed a series of IR spectroscopy and IR-SEC experiments to measure the hydricity and pKa values of pertinent reaction intermediates. Comparison of the thermochemical data for these intermediates in both MeCN/H2O (95:5) and buffered aqueous solution is included. Extensive reports on the thermochemical properties of H2evolving and CO2-reducing electrocatalysts have appeared for systems studied in MeCN.35 However, data in aqueous media are limited,36 despite the intense interest in developing electrocatalysts for CO2 reduction to C−H bond-containing products that operate in water. Catalyst 1− is effective for selective formate production in water, and, in this context, we have probed the pKa values for 1− and the corresponding reduced hydride intermediates in water. This has allowed us to compare thermochemical data in MeCN and water for the same catalyst and provide a benchmark to the extensive data available in the literature for MeCN. Acidity Measurement in MeCN for [HFe4N(CO)12]− (H− 1) . If (H-1)− will transfer H− to CO2, then it is the nature of this intermediate that will determine whether H2 or formate production can occur in solutions of 1− and CO2 in the presence of protons. To determine the pKa of (H-1)−, the following equilibrium was investigated:

where pK a = pKeq + pK a(PhCO2 H)

(9) 40

Similar to methods reported by Sarker and Bruno, as well as others, the equilibrium constant (Keq) was determined from the intensities of CO stretches. Intensities of IR transitions 7145

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ACS Catalysis collected in absorption mode obey Beer’s law and can be used as suitable probes of solution concentrations.41 As an example, for 1−, absorption versus concentration of 1− provides a linear relationship (Figure 9, left). Other related compounds, 12− and

Figure 10. IR spectra collected before (black trace) and after (red trace) application of −1.2 V during CPE experiments on aqueous buffered solutions of 0.3 mM 1−: (left) pH 9 and (right) pH 13.

10−13, but that the reaction does not go to completion. From this experiment, we estimate that the pKa of (H-1)− falls within one unit of 13 (Table 3). These measurements indicate that protonation of 12− to afford (H-1)− is facile in aqueous solution, even at pH 13. Hydricity in MeCN, for [HFe4N(CO)12]− (H-1)−. Hydride transfer is an important step in both H2 formation and CO2 reduction,45 because an appropriate catalyst for CO2 reduction must be hydridic enough to facilitate C−H bond formation. Hydricity is defined as the free energy of heterolytic metalhydride bond cleavage as described generally in eq 10:

Figure 9. (Left) IR absorption spectra for different concentrations of 1− in MeCN. Inset shows the absorption (Abs) at 1989 cm−1 vs [1−]. Overlaid pink line is the linear fit. (Right) Normalized difference spectra showing: 1− turning into (H-1)− (blue trace); 1− turning into 12− (black trace); the equilibrium between 12− + PhCOOH and (H1)− + PhCOOH− (eq 6) (red trace); and simulated spectrum obtained from the summation of 0.7(black) + 0.3(blue) (pink trace).

H-1−, are formed upon the reduction of 1− in dry acetonitrile and the reduction of 1− in MeCN:H2O (95:5), respectively. Therefore, we normalized the final difference spectra acquired under both conditions to the same negative absorption for 1− to indicate that the same amount of 1− was used to produce 12− and (H-1)−. Similarly, we normalized the difference absorption spectra acquired upon the reduction of 1− in dry acetonitrile in the presence of benzoic acid (Figure 9, right red spectrum). Quantification of the ratio of 12− and (H-1)− in this final spectrum was done by finding a ratio of two spectra related to 12− and (H-1)−, represented by the black and blue spectra, that matches the red spectrum the best. The pink spectrum is the summation of 70% 1− and 30% (H-1)−. This simulation afforded us Keq for the equilibrium shown in eq 6. Using the thermochemical cycle shown in eqs 6−8, and eq 9, the pKa for (H-1)− was then calculated as 20.3 (shown in Table 3; see the Supporting Information for details of the calculation).

M−H ⇌ M+ + H−

ΔG H° −

(10)

There are several methods commonly used to measure the equilibrium shown in eq 14.42 One approach uses the thermochemical cycle shown in eqs 11−14, where the reactivity of 1 with H2 should be interrogated to establish the equilibrium shown in eq 11. In eq 13, ΔG°H2 is the standard free energy for heterolytic H2 cleavage, and is equal to 76 kcal mol−1 in MeCN.46 Compound 1 has not been isolated and so we could not study its reactivity with H2. Under catalytic conditions, 1 is quickly converted to 1−; therefore, long-term stability is not a problem during catalysis.

Table 3. Thermochemical Properties of (H-1)−, H2, and Formate in MeCN and H2O compound

pKa (MeCN)

pKa (H2O)

ΔGH°− (MeCN) (kcal mol−1)

ΔGH°− (H2O) (kcal mol−1)

(H-1)− H2

20.3 55.5 (ref 42)

13 ± 1 25.1 (ref 43)

49 76 (ref 42)

15.5 34.2 (ref 43)

44 (ref 44)

24.1 (ref 43)

HCOO−

where ΔG H° − = ΔG° (eq 11) + ΔG° (eq 12) + ΔG H° 2 = 1.37pKeq − 1.37pK acid + 76 kcal mol−1

−

−

Acidity in Water for [HFe4N(CO)12] (H-1) . IR changes were investigated after reducing 1− to 12− at −1.2 V in aqueous solutions buffered at various pH levels under an H2 atmosphere in a CPE experiment (eq 6). In solutions with pH ≤12, IR spectra before and after passing 1−2 equiv of charge were identical, and this indicates that any reduced cluster generated is reacting with protons to afford H2 and regenerate 1− (Figure 10, left). At pH 13, we observed a new absorption at ∼1940 cm−1, which is consistent with the formation of (H-1)− (Figure 10, right). The presence of (H-1)− in this spectrum indicates that 12− is basic enough to react with H+ with concentration of

(15)

Since we were not able to measure Keq directly via a reaction, as shown in eq 11, we used an alternative approach. We studied the reactivity of (H-1)− with acids of various strengths and acquired two limits for Keq: a strong acid that converts (H-1)− completely to H2 gas provides us with the upper limit for the value of Keq, and a weaker acid, which does not significantly change the concentration of (H-1)−, provides us the lower limit for its value. Using eqs 11−15, the hydride donor ability of (H1)− was then estimated. Using IR-SEC in 0.1 M Bu4NPF6 MeCN solution, (H-1)− was generated quantitatively by application of −1.4 V potential 7146

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ACS Catalysis to a solution of 1− and 1 equiv of an organic acid under 1 atm H2, both in the presence and absence of H2O (4%). In both cases, when benzoic acid (pKa = 20.7 MeCN)39 was used, no H2 was generated.47 The addition of 10 equiv of benzoic acid also afforded no H2. Therefore, we can assume that Keq for the reaction with benzoic acid is 10, and then the hydricity can be estimated as ≥49 kcal mol−1 (see Table 3). Taken together, these experiments indicate that ΔG°H− for (H1)− is 49 kcal mol−1. The detailed calculation is shown in the Supporting Information. The hydride donor ability of the formate anion in MeCN has been estimated as 44 kcal mol−1.44 This comparison indicates that transfer of H− from (H-1)− to CO2 is thermodynamically unfavorable by 5 kcal mol−1 in MeCN. The important effect of the added water on facilitating C−H bond formation is detailed below. Homolytic Bond Dissociation Free Energy (BDFE) in MeCN, for [HFe4N(CO)12]− (H-1)−. We also estimated the homolytic bond dissociation free energy (BDFE) for (H-1)− of the metal hydride, as described generally in eq 16, and specifically for (H-1)− in eq 17. Commonly, the value of this free energy is calculated using eq 18.48 The constants (expressed in units of kcal mol−1) shown in eq 16 are appropriate when MeCN is used as a solvent and potentials are referenced to the ferrocenium/ferrocene (Fc+/0) couple. M−H ⇌ M + H•

where ΔG H° − = ΔG°(eq 19) + ΔG H° 2 = 1.37pKeq + 34.2 kcal mol−1 Keq =

BDFE (17)

BDFE = 1.37pK a + 23.06E°(1− /12 − ) + 53.6 kcal mol−1 (18) −

⎛ [(H‐1)− ] ⎞ → pKeq = log⎜ ⎟ − pH ⎝ [1] ⎠ [(H‐1) ][H ] [1]pH −

2

+

(23)

1−

−

(22)

The equilibrium constant for eq 19 (Keq) can be calculated using eq 23. In this equation, [1], [(H-1)−], and [H+] are the concentration of each species and pH2 is the pressure of hydrogen gas (pH2 = 1 atm). We have already described that (H-1)− is observed in solution at pH 13, following CPE experiments performed under 1 atm of H2 gas (Figure 10, right). Therefore, we can assume that [(H-1)−]/[1] > 0.1 at pH 13. Then, using eq 23, pKeq ≥ −14. In addition, we were not able to observe any (H-1)− at lower pH conditions (pH 5− 12, Figure 10, left), and, from this observation, we can estimate that [(H-1)−]/[1] < 0.1 at pH 12, and then the pKeq can be estimated as being less than or equal to −13. From these combined results, we concluded that −14 ≤ pKeq ≤ −13 for eq 19. Using eq 22, we can then calculate that, for (H-1)−, 15 kcal mol−1 ≤ ΔG°H− ≤ 16 kcal mol−1. Using a value of 24.1 kcal mol−1 for hydride donor ability of formate,43 our measurements predict that the reaction of (H-1)− with CO2 in water should have a driving force of 8.6 kcal mol−1 for C−H bond formation. Homolytic Bond Dissociation Free Energy (BDFE) in Aqueous Solution for [HFe4N(CO)12]− (H-1)−. BDFE values in water could not be calculated because we observed an electrocatalytic reduction of protons to H2, and this precludes the observation of reversible electrochemical processes for the 1−/2− couple. These are necessary to determine the BDFE values (eq 18). Thermochemical Measurements Explain the Observed CO2 Reduction Reactivity. Despite the importance of the CO2 reduction reaction in aqueous solution as a potential fuel-forming process,51 limited reports on thermochemical measurements of any metal hydride, and especially on functional formate-producing catalysts under these conditions, have appeared. 52 Creutz, Muckerman, and co-workers compared the hydricity values for two Ru complexes in both water and MeCN,46a and found a general compression of the

(16)

BDFE

[HFe4N(CO)12 ]− ⇌ [Fe4N(CO)12 ]− + H• (H ‐ 1)−

observed with organic solvents, but instead of using organic acids with varied pKa values, the solution acidity can be readily changed by adjusting the pH using aqueous buffers. Using the thermochemical cycle outlined in eqs 19−21, the values for ΔG°H− of (H-1)− in water were calculated. Two previously reported estimates of ΔG°H2 in water were made by Pearson, and by Kelly and Rosseinsky, and these are different by 13 kcal mol−1.50 Here, we use a value of 34.2 kcal mol−1 which was reported recently by Appel and co-workers,43 and is the modification of the reported number by Kelly and Rosseinsky.50a This value is obtained from a rigorous treatment of experimental data.43

2−

Hydricity values measured for (H-1) , and E°(1 /1 ) vs Fc+/0 (vide supra) were used to obtain the BDFEs for (H-1)−. According to eq 18, the BDFE for (H-1)− is 47 kcal mol−1, and this is 2 kcal mol−1 lower than ΔG°H− for (H-1)−. As a comparison, the BDFE for H2 in acetonitrile is 103.6 kcal mol−1,49 and this is significantly larger than its hydricity value, ΔG°H− = 76 kcal mol−1. At first glance, it appears that the loss of H• from (H-1)− in MeCN is more favorable, by 2 kcal mol−1, than the loss of H− (47 kcal mol−1 vs 49 kcal mol−1, respectively). However, comparison of the overall reactions of interest must be considered: H2 formation is described here as an example. The free energy for the loss of H2 in the reaction 2(H-1)− → 21−+ H2, is −9 kcal mol−1, and the free energy for the transfer of hydride to a proton in the reaction (H-1)− + H+ → 1− + H2, is −27 kcal mol−1. Therefore, although the BDFE of (H-1)− is small, a homolytic pathway is less favorable than a heterolytic pathway in the production of H2 by (H-1)−. It should also be noted that we do not observe the release of H2 gas in IR-SEC experiments where we generate (H-1) − quantitatively. This implies that bimolecular release of H2 to form 1− is not favorable. Taken together, these results convince us that hydride transfer to substrates such as H+ or CO2 is a reasonable reaction pathway involving (H-1)−. Hydricity in Water, for [HFe4N(CO)12]− (H-1)−. To measure hydricity in water, the principle is the same as that 7147

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(THF): 1952 (w, 12−), 1909 (sh, 12−), 1896 (s, 12−), 1856 (sh, 12−), 1832 (sh, 12−), 1896 (s, (H-1)−), 1885 (s, (H-1)−), and 1854 (sh, (H-1)−) cm−1. We were unable to obtain elemental analysis because the crystals will decompose when warmed to room temperature. Reaction of (12−)/(H-1)− with CO2. A sample of 12−/(H-1)− (1 mg, 0.5 mmol) cooled in an acetone/dry ice bath under a CO2 atmosphere was prepared and then chilled MeCN (4 mL) was added. The temperature of the reaction was then raised to room temperature over 1 h before the MeCN was removed under vacuum. Four equivalents (4 equiv) of KCl were added and 1H NMR analysis of the water-soluble portion of the resulting precipitate revealed formate in 25% ± 5% yield, based on (H-1)− (see Figure S19). Formate was quantified using proton NMR spectroscopy, as described in the section titled “Electrochemical Measurements”. Electrochemical Measurements. Cyclic voltammograms were recorded under a nitrogen gas (N2) (99.998%, Praxair) atmosphere using an electrochemical analyzer (CH Instruments, Model 620D or Model 1100B), a glassy carbon working electrode (CH Instruments) with a nominal surface area of 0.0707 cm2), and a platinum wire auxiliary electrode. The glassy carbon working electrode was polished on a felt pad with alumina paste (0.05 μm, BASi), sonicated in deionized water, rinsed with MeOH, and dried with a Kimwipe prior to each experiment. As the reference electrode, a Ag/AgCl(sat.)/1 M KCl or a nonaqueous Ag/AgNO3 (0.001 M) was used for aqueous and nonaqueous measurements, respectively. All reported potentials are referenced to the SCE couple, and were determined using ferrocene (Aldrich) as an internal standard, where E1/2(Fc+1/0) is +0.159 V vs SCE in water and +0.40 V vs SCE in acetonitrile.22b,53 Nonaqueous electrolyte solutions were stored over 3 Å molecular sieves that had been activated by heating under vacuum at 200 °C for at least 72 h. Milli-Q water (18 MΩ) was used for measurements performed in aqueous solution. Buffer solutions were all 0.1 M and phosphate buffer was used for pH 6−8, acetate for pH 5, borate for pH 9, and sodium hydroxide−potassium chloride for pH 13. Reagents for buffer preparation were purchased from EMD, VWR, and Sigma, and were used as received. In all cases, Cyclic voltammetry (CV) sweeps were initiated at the open circuit potential and recorded in quiescent solution. No iR compensation was used for electrochemical measurements. Controlled Potential Electrolysis (CPE). CPE experiments were performed in a gas-tight glass cell (working electrode compartment volume of 60 mL) under 1 atm of static N2 (Praxair, 99.998%) or carbon dioxide collected from dry ice with a stirred solution. The counter electrode compartment was separated from the working electrode compartment by a glass frit of medium porosity. In a typical experiment, 25 mL of degassed electrolyte solution was used in the working electrode compartment. Gas samples were injected directly into a gas chromatography (GC) system (see the section titled “Other Physical Measurements”). The working electrode was a glassy carbon plate with an area of 2.5 cm2 (Tokai Carbon), while the counter electrode was a coiled Pt wire ∼30 cm in length (BASi). The reference electrodes employed for CPE experiments were of similar design to those used for CV measurements. Between CPE experiments, the glass cell, the stir bar, the working electrode, and the counter electrode were cleaned via sonication in 5% (v/v) nitric acid for 5 min, rinsed with deionized (DI) water, sonicated in DI water for 5 min, rinsed with DI water, and then sonicated in acetone for 5 min,

hydricity range, and a decrease in hydricity values, in H2O, compared to MeCN. We measured ΔGH°− for (H-1)− as 15.5 kcal mol−1 in water, and 49 kcal mol−1 in MeCN. These results indicate that the driving force for C−H bond formation by (H-1)− is favorable by 8.5 (in H2O) and unfavorable by 5 kcal mol−1 (in MeCN), and this also corroborated our experimental observations where we observed greater selectivity for formate production over H2 formation upon moving from pure MeCN, to MeCN/H2O (95:5), and then even better selectivity in buffered aqueous solution. The rate of formation of formate in CPE experiments in aqueous solution was also higher than in MeCN. Important catalyst design elements are illustrated here. The hydricity of the catalytic intermediate (H-1)− and the driving force for C−H bond formation with CO2 are relatively modest, and the reaction outcomes are highly dependent on the reaction medium. This situation has enabled selective C−H bond formation in water with CO2, i.e., when the hydricity of the H− donor is ∼8 kcal mol−1 more favorable than HCOO−, C−H bond formation with CO2 will result.

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CONCLUSION We have described an Earth-abundant metal catalyst [Fe4N(CO)12]− (1−) that is selective for the reduction of CO2 to formate in aqueous solution. The fastest rates of reaction observed in this study were observed at pH 7, and under these conditions, controlled potential electrolysis (CPE) experiments were operated at 4 mA cm−2 for over 24 h and formate was consistently generated with 95% Faradaic efficiency (background H2 evolution results from reaction at the glassy carbon electrode and not the catalyst itself). The high stability of 1−, combined with high selectivity toward formate formation over CO, H2, oxalate, methanol, or formaldehyde is unusual among transition-metal electrocatalysts (either homogeneous or heterogeneous), or organocatalyzed electroreduction reactions.10−18 We have combined these catalytic results with thermochemical data that rationalize the observed selectivity for formate by [Fe4N(CO)12]− (1−). Ongoing work targets new catalysts, collaborative mechanistic investigations using density functional theory (DFT), and other strategies for selective formation of multiple C−H bonds with CO2, based on the insights gained from this work.

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EXPERIMENTAL SECTION Reagents and Procedures. All manipulations were carried out using standard Schlenk or glovebox techniques under a nitrogen gas (N2) atmosphere. Unless otherwise noted, solvents were deoxygenated and dried by thorough sparging with Ar gas, followed by passage through an activated alumina column. Diglyme was purchased in a Sureseal bottle, degassed, and dried over Na, distilled under reduced pressure, and stored over activated 3 Å molecular sieves. Deuterated solvents were purchased from Cambridge Isotopes Laboratories, Inc. and were degassed and stored over activated 3 Å molecular sieves prior to use. [(dyglyme)2Na][Fe4NCO)12] (1−) was prepared using previously reported methods.20 [Co(C5Me5)2]3[Fe4N(CO)12][HFe4N(CO)12] (12−)/(H-1)−. One equivalent (1 equiv) of Cp2*Co (19 mg, 58 mmol) was added to (Na-1), (50 mg, 58 mmol) dissolved in 5 mL of THF and the solution was stirred for 1 day. Crystals suitable for a singlecrystal X-ray diffraction experiment were obtained by diffusion of ether into the THF solution over 4 months at −25 °C. IR 7148

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anisotropically. Hydrogen atoms, where added, were assigned to ideal positions and refined using a riding model with an isotropic thermal parameter 1.2 times greater than that of the attached carbon atom (1.5 times greater than that for methyl hydrogens). The possible positions of the hydride in 12−/(H-1)− are either bridging the Fe2−Fe2a bond (the butterfly hinge) at 50% occupancy, or bridging the two equivalent Fe2−Fe3 bonds (the butterfly wing edges), each at 25% occupancy, because of a crystallographically imposed mirror plane. We investigated these two options by calculating the initial hydride position using DFIX and then freely refining the thermal parameter on the H atom. In the Fe2−Fe2a butterfly hinge position, the thermal parameter was refined to U = 0.56, which indicated that this solution is unlikely to be correct. When the H atom was included bridging the Fe2−Fe3 bond, free refinement of the H atom afforded U = 0.065. Thus, we believe that the correct solution to the data is one in which the hydride bridges Fe2− Fe3 along the butterfly wing. Other Physical Measurements. Electrospray ionization− mass spectrometry (ESI-MS) was performed using an Agilent Model 1100LC/MSD G1956b system that was equipped with single quadruple at a cone voltage of 20 V. The sample solution was directly injected via spray nozzle with a syringe pump at a speed of 0.1 mL/min.

rinsed with DI water, and allowed to dry in an oven before use. The glassy carbon plate had an additional initial step of being thoroughly sanded on all surfaces with 300 grit SiC paper and then 600 grit SiC paper and rinsed with water prior to sonication steps. Quantification of headspace gases and solution-phase products were performed separately. At the end of an electrolysis, a gaseous sample (0.1 mL) was drawn from the headspace, using a gas-tight syringe (Vici), and injected into a gas chromatography−thermal conductivity detection (GCTCD) system (Model Varian 3800 GC coupled with a TCD detector and a Carboxen 1010 PLOT fused silica column (30 m × 0.53 mm) (Supelco) using N2 (99.999%, Praxair) as the carrier gas. H2 concentration was determined using a previously prepared working curve (see Figure S8). Analysis of the liquid phase was achieved by removing 0.5 mL of sample from the CPE experiment. 1H NMR analysis was performed using a 600 MHz Varian spectrometer on each sample. The same capillary containing a spike of dimethylformamide (DMF) in C6D6 was included in each experiment. In a separate experiment, a plot of [NaCHOO] vs the ratio of NMR integrals of DMF to formate was constructed from a series of NMR spectra collected of aqueous solutions of sodium formate. This plot was used to determine the concentration of formate produced in CPE experiments (see Figure S6). Turnover numbers (TONs) were calculated based on the total amount of product (formate or hydrogen gas) divided by the concentration of catalyst used in CPE experiments. Spectroscopic and Spectroelectrochemical Measurements. Infrared spectra were recorded on a Bruker Alpha Infrared spectrometer (2 cm−1 resolution). Infrared spectroelectrochemical (IR-SEC) measurements were performed using an optically transparent, thin-layer solution infrared (IR) cell fabricated by Prof. Hartl at the University of Reading in the United Kingdom, as has been described previously.54 The IRSEC cell was gas-tight and contained a masked Au-minigrid working electrode (32 wires/cm), a Pt-gauze auxiliary electrode, and an Ag-wire pseudo-reference electrode, and CaF2 windows. In each experiment, electrochemical reduction of the species of interest was monitored by IR spectroscopy for a period of 2−15 min. Diffusion and mixing of the redox products, generated at the working and auxiliary electrodes in the IR cell, were reasonably suppressed within the total experimental time (typically 10−15 min for one complete measurement). Degassed and air-free samples were loaded directly into the IR-SEC cell, using a Luer-lock syringe. X-ray Structure Determination. X-ray diffraction (XRD) studies were carried out on a Bruker SMART APEX Duo diffractometer that was equipped with a CCD detector.55 Measurements were carried out at −175 °C, using Mo Kα (λ = 0.71073 Å) radiation. Crystals were mounted on a glass capillary or Kaptan Loop with Paratone-N oil. Initial lattice parameters were obtained from a least-squares analysis of more than 100 centered reflections; these parameters were later refined against all data. Data were integrated and corrected for Lorentz polarization effects using SAINT and were corrected for absorption effects using SADABS2.3. Space group assignments were based upon systematic absences, E statistics, and successful refinement of the structures. Structures were solved by direct methods with the aid of successive difference Fourier maps and were refined against all data using the SHELXTL 2014/7 software package. Thermal parameters for all non-hydrogen atoms were refined

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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/acscatal.5b01708. CPE results, calibration curves for formate quantification, plots of CV measurements, and calculation for thermochemical values (PDF) Crystallographic data of C84H91Co3Fe8N2O24 (CIF)

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Author Contributions

All authors have given approval to the final version of the manuscript. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by UC Davis and by the National Science Foundation through the CAREER Program (No. CHE1055417). L.A.B. is an Alfred P. Sloan Foundation Fellow.

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REFERENCES

(1) Lüthi, D.; Le Floch, M.; Bereiter, B.; Blunier, T.; Barnola, J.-M.; Siegenthaler, U.; Raynaud, D.; Jouzel, J.; Fischer, H.; Kawamura, K.; Stocker, T. F. Nature 2008, 453, 379−382. (2) (a) Lewis, N. S.; Nocera, D. G. Proc. Natl. Acad. Sci. U. S. A. 2006, 104, 15729−15735. (b) Olah, G. A.; Prakash, G. K. S.; Goeppert, A. J. Am. Chem. Soc. 2011, 133, 12881−12898. (c) Goeppert, A.; Czaun, M.; Jones, J. P.; Surya Prakash, G. K.; Olah, G. A. Chem. Soc. Rev. 2014, 43, 7995−8048. (3) Halmann, M. M.; Steinberg, M. Greenhouse Gas Carbon Dioxide Mitigation: Science and Technology; Lewis Publishers: Boca Raton, FL, 1999. (4) See, for example: (a) Riplinger, C.; Sampson, M. D.; Ritzmann, A. M.; Kubiak, C. P.; Carter, E. A. J. Am. Chem. Soc. 2014, 136, 7149

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ACS Catalysis 16285−16298. (b) Sampson, M. D.; Nguyen, A. D.; Grice, K. A.; Moore, C. E.; Rheingold, A. L.; Kubiak, C. P. J. Am. Chem. Soc. 2014, 136, 5460−5471. (c) Costentin, C.; Passard, G.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2014, 136, 11821−11829. (d) Fisher, B. J.; Eisenberg, R. J. Am. Chem. Soc. 1980, 102, 7361−7363. (e) Bruce, M. R. M.; Megehee, E.; Sullivan, B. P.; Thorp, H.; O'Toole, T. R.; Downard, A.; Meyer, T. J. Organometallics 1988, 7, 238−240. (f) Collin, J. P.; Jouaiti, A.; Sauvage, J. P. Inorg. Chem. 1988, 27, 1986−1990. (5) See, for example: (a) Hori, Y. In Modern Aspects of Electrochemistry, Vol. 42; Vayenas, C. G., White, R. E., GamboaAldeco, M. E., Eds.; Springer: New York, 2008; pp 89−190. (b) Peterson, A. A.; Abild-Pedersen, F.; Studt, F.; Rossmeisl, J.; Nørskov, J. K. Energy Environ. Sci. 2010, 3, 1311−1315. (6) See, for example: (a) Bhugun, I.; Lexa, D.; Savéant, J. M. J. Am. Chem. Soc. 1996, 118, 1769−1776. (b) Caix, C.; Chardon-Noblat, S.; Deronzier, A. J. Electroanal. Chem. 1997, 434, 163−170. (c) Kushi, Y.; Nagao, H.; Nishioka, T.; Isobe, K.; Tanaka, K. Chem. Lett. 1994, 2175−2178. (d) Ishida, H.; Tanaka, H.; Tanaka, K.; Tanaka, T. J. J. Chem. Soc., Chem. Commun. 1987, 131−132. (e) Ali, M. M.; Sato, H.; Ali, M. M.; Sato, H.; Mizukawa, T.; Tsuge, K.; Haga, M.; Tanaka, K. Chem. Commun. 1998, 2, 249−250. (7) See, for example: (a) Kim, D.; Resasco, J.; Yu, Y.; Asiri, A. M.; Yang, P. Nat. Commun. 2014, 5, 4948−4956. (b) Rasul, S.; Anjum, D. H.; Jedidi, A.; Minenkov, Y.; Cavallo, L.; Takanabe, K. Angew. Chem., Int. Ed. 2015, 54, 2146−2150. (c) Zhu, W.i; Zhang, Y.-J.; Zhang, H.; Lv, H.; Li, Q.; Michalsky, R.; Peterson, A. A.; Sun, S. J. Am. Chem. Soc. 2014, 136, 16132−16235. (d) Li, C.; Kanan, M. W. J. Am. Chem. Soc. 2012, 134, 7231−7234. (e) Sen, S.; Liu, D.; Palmore, G. T. R. ACS Catal. 2014, 4, 3091−3095. (f) Huan, T. N.; Andreiadis, E. S.; Heidkamp, J.; Simon, P.; Derat, E.; Cobo, S.; Royal, G.; Bergmann, A.; Strasser, P.; Dau, H.; Artero, V.; Fontecave, M. J. Mater. Chem. A 2015, 3, 3901−3907. (g) Reske, R.; Mistry, H.; Behafarid, F.; Roldan Cuenya, B.; Strasser, P. J. Am. Chem. Soc. 2014, 136, 6978−6986. (h) Zhang, Y.-J.; Peterson, A. A. Phys. Chem. Chem. Phys. 2015, 17, 4505−4515. (8) For example: (a) Lim, C.-H.; Holder, A. M.; Hynes, J. T.; Musgrave, C. B. J. Am. Chem. Soc. 2014, 136, 16081−16095. (b) Xiang, D.; Magana, D.; Dyer, R. B. J. Am. Chem. Soc. 2014, 136, 14007− 14010. (c) Boston, D. J.; Xu, C.; Armstrong, Da. W.; MacDonnell, F. M. J. Am. Chem. Soc. 2013, 135, 16252−16255. (9) (a) Manthiram, K.; Beberwyck, B. J.; Alivisatos, A. P. J. Am. Chem. Soc. 2014, 136, 13319−13325. (b) Nakata, K.; Ozaki, T.; Terashima, C.; Fujishima, A.; Einaga, Y. Angew. Chem., Int. Ed. 2014, 53, 871−875. (10) (a) Kang, P.; Cheng, C.; Chen, Z.; Schauer, C. K.; Meyer, T. J.; Brookhart, M. J. Am. Chem. Soc. 2012, 134, 5500−5503. (b) Kang, P.; Meyer, T. J.; Brookhart, M. Chem. Sci. 2013, 4, 3497−3502. (11) Kang, P.; Zhang, S.; Meyer, T. J.; Brookhart, M. Angew. Chem., Int. Ed. 2014, 53, 8709−8713. (12) (a) Azuma, M.; Hashimoto, K.; Hiramoto, M.; Watanabe, M.; Sakata, T. J. Electroanal. Chem. Interfacial Electrochem. 1989, 260, 441− 445. (b) Ikeda, S.; Takagi, T.; Ito, K. Bull. Chem. Soc. Jpn. 1987, 60, 2517−2522. (13) Lee, C. H.; Kanan, M. W. ACS Catal. 2015, 5, 465−469. (14) Podlovchenko, B. I.; Kolyadko, E. A.; Lu, S. J. Electroanal. Chem. 1994, 373, 185−187. (15) Kortlever, R.; Balemans, C.; Kwon, Y.; Koper, M. T. M. Catal. Today 2015, 244, 58−62. (16) Min, X.; Kanan, M. W. J. Am. Chem. Soc. 2015, 137, 4701−4708. (17) See, for example: (a) Hori, Y.; Kikuchi, K.; Suzuki, S. Chem. Lett. 1985, 14, 1695−1698. (b) Chen, Y.; Kanan, M. W. J. Am. Chem. Soc. 2012, 134, 1986−1989. (c) Machunda, R. L.; Ju, H.; Lee, J. Curr. Appl. Phys. 2011, 11, 986−988. (d) Agarwal, A. S.; Zhai, Y.; Hill, D.; Sridhar, N. ChemSusChem 2011, 4, 1301−1310. (18) Zhang, S.; Kang, P.; Meyer, T. J. J. Am. Chem. Soc. 2014, 136, 1734−1737. (19) Rail, M. D.; Berben, L. A. J. Am. Chem. Soc. 2011, 133, 18577− 18579.

(20) Nguyen, A. D.; Rail, M. D.; Shanmugam, M.; Fettinger, J. C.; Berben, L. A. Inorg. Chem. 2013, 52, 12847−12854. (21) (a) Felton, G. A. N.; Vannucci, A. K.; Okumura, N.; Lockett, L. T.; Evans, D. H.; Glass, R. S.; Lichtenberger, D. L. Organometallics 2008, 27, 4671−4679. (b) At 5% water concentration, homoconjugation is expected for water, although we could not find a value in the literature. (22) (a) Randles, J. E. B. Trans. Faraday Soc. 1948, 44, 327−338. (b) Bard, A. J.; Faulkner, L. R.; Swain, E.; Robey, C. Electrochemical Methods: Fundamentals and Applications; 2nd Edition; John Wiley and Sons: New York, 2001. (23) Costentin, C.; Drouet, S.; Robert, M.; Savéant, J.-M. Science 2012, 338, 90−94. (24) (a) Savéant, J.-M.; Vianello, E. Electrochim. Acta 1965, 10, 905− 920. (b) Saveánt, J. M.; Vianello, E. Electrochim. Acta 1967, 12, 629− 646. (c) Savéant, J.-M. Elements of Molecular and Biomolecular Electrochemistry: An Electrochemical Approach to Electron Transfer Chemistry; Wiley−Interscience: Hoboken, NJ, 2006. (d) Costentin, C.; Savéant, J.-M. ChemElectroChem 2014, 1, 1226−1236. (e) Artero, V.; Savéant, J.-M. Energy Environ. Sci. 2014, 7, 3808−3814. (25) (a) Gennaro, A.; Isse, A. A.; Vianello, E. J. Electroanal. Chem. Interfacial Electrochem. 1990, 289, 203−215. (b) Fujita, E.; Szalda, D. J.; Creutz, C.; Sutin, N. J. Am. Chem. Soc. 1988, 110, 4870−4871. (26) See Figure S16 in the Supporting Information for an alternative method to calculate the rate constants. (27) Carroll, J. J.; Slupsky, J. D.; Mather, A. E. J. Phys. Chem. Ref. Data 1991, 20, 1201−1209. (28) See the Supporting Information for an explanation of the procedure used to obtain difference absorption (DAbs) and normalized spectra. (29) (a) Smieja, J. M.; Kubiak, C. P. Inorg. Chem. 2010, 49, 9283− 9289. (b) Cheah, M. H.; Borg, S. J.; Bondin, M. I.; Best, S. P. Inorg. Chem. 2004, 43, 5635−5644. (c) Brooksby, P. A.; Duffy, N. W.; Mcquillan, A. J.; Robinson, B. H.; Simpson, J. J. Chem. Soc., Dalton Trans. 1998, 3, 2855−2860. (d) Golovin, M. N.; Weaver, M. J. Inorg. Chim. Acta 1988, 142, 177−179. (e) Bullock, J. P.; Palazzotto, M. C.; Mann, K. R. Inorg. Chem. 1990, 29, 4413−4421. (f) Best, S. P.; Clark, R. J. H.; Deeming, A. J.; McQueen, R. C. S.; Powell, N. I.; Acuña, C.; Arce, A. J.; De Sanctis, Y. J. Chem. Soc., Dalton Trans. 1991, 1111− 1115. (30) Tachikawa, M.; Stein, J.; Muetterties, E. L.; Teller, R. G.; Beno, M. A.; Gebert, E.; Williams, J. M. J. Am. Chem. Soc. 1980, 102, 6648− 6649. (31) Blaine, C. A.; Ellis, J. E.; Mann, K. R. Inorg. Chem. 1995, 34, 1552−1561. (32) Cheng, S. C.; Blaine, C. A.; Hill, M. G.; Mann, K. R. Inorg. Chem. 1996, 35, 7704−7708. (33) See, for example: (a) Hori, Y.; Murata, A.; Kikuchi, K.; Suzuki, S. J. Chem. Soc., Chem. Commun. 1987, 728−729. (b) Kedzierzawski, P.; Augustynski, J. J. Electrochem. Soc. 1994, 141, L58−L60. (c) Ohmori, T.; Nakayama, a.; Mametsuka, H.; Suzuki, E. J. Electroanal. Chem. 2001, 514, 51−55. (d) Stevens, G. B.; Reda, T.; Raguse, B. J. Electroanal. Chem. 2002, 526, 125−133. (34) See, for example: (a) Haynes, L. V.; Sawyer, D. T. Anal. Chem. 1967, 39, 332−338. (b) Ohta, K.; Kawamoto, M.; Mizuno, T.; Lowy, D. A. J. Appl. Electrochem. 1998, 28, 717−724. (c) Kaneco, S.; Iiba, K.; Ohta, K.; Mizuno, T.; Saji, A. J. Electroanal. Chem. 1998, 441, 215− 220. (35) See, for example: (a) Jeletic, M. S.; Mock, M. T.; Appel, A. M.; Linehan, J. C. J. Am. Chem. Soc. 2013, 135, 11533−11536. (b) Wiese, S.; Kilgore, U. J.; Dubois, D. L.; Bullock, R. M. ACS Catal. 2012, 2, 720−727. (c) Kilgore, U. J.; Roberts, J. A. S.; Pool, D. H.; Appel, A. M.; Stewart, M. P.; Dubois, M. R.; Dougherty, W. G.; Kassel, W. S.; Bullock, R. M.; Dubois, D. L. J. Am. Chem. Soc. 2011, 133, 5861−5872. (36) (a) Fu, X.; Wayland, B. B. J. Am. Chem. Soc. 2004, 126, 2623− 2631. (b) Bhagan, S.; Wayland, B. B. Inorg. Chem. 2011, 50, 11011− 11020. (c) Fu, X.; Li, S.; Wayland, B. B.; Uni, V.; Pennsyl, V. Inorg. Chem. 2006, 45, 9884−9889. (d) Chou, M.; Creutz, C.; Mahajan, D.; Sutin, N.; Zipp, A. P. Inorg. Chem. 1982, 21, 3989−3997. (e) Creutz, 7150

DOI: 10.1021/acscatal.5b01708 ACS Catal. 2015, 5, 7140−7151

Research Article

ACS Catalysis C.; Chou, M. H. J. Am. Chem. Soc. 2007, 129, 10108−10109. (f) Creutz, C.; Chou, M. H. J. Am. Chem. Soc. 2009, 131, 2794−2795. (g) Schwarz, H. A.; Creutz, C. Inorg. Chem. 1983, 22, 707−713. (37) (a) Pelmenschikov, V.; Guo, Y.; Wang, H.; Cramer, S. P.; Case, D. A. Faraday Discuss. 2011, 148, 409−420. (b) Franke, O.; Wiesler, B. E.; Lehnert, N.; Tuczek, F. Z. Z. Anorg. Allg. Chem. 2002, 628, 2395− 2402. (c) Brunet, J.-J.; Chauvin, R.; Chiffre, J.; Huguet, S.; Leglaye, P. J. Organomet. Chem. 1998, 566, 117−123. (d) Kaesz, H. D.; Saillant, R. B. Chem. Rev. 1971, 72, 231−281. (38) This conclusion, whether correct or not, has no consequences on the following discussion. (39) Izutsu, K. Acid−Base Dissociation Constants in Dipolar Aprotic Solvents; IUPAC Chemical Data Series, No. 35; Blackwell Science: Oxford, U.K., 1990. (40) Sarker, N.; Bruno, J. W. J. Am. Chem. Soc. 1999, 121, 2174− 2180. (41) See, for example: (a) Moore, E. J.; Sullivan, J. M.; Norton, J. R. J. Am. Chem. Soc. 1986, 108, 2257−2263. (b) Kristjansdottir, S. S.; Moody, A. E.; Weberg, R. T.; Norton, J. R. Organometallics 1988, 7, 1983−1987. (c) Jordan, R. F.; Norton, J. R. J. Am. Chem. Soc. 1982, 104, 1255−1263. (42) Curtis, C. J.; Miedaner, A.; Ellis, W. W.; Dubois, D. L. J. Am. Chem. Soc. 2002, 124, 1918−1925. (43) Connelly, S. J.; Wiedner, E. S.; Appel, A. M. Dalton Trans. 2015, 44, 5933−5938. (44) (a) Price, A. J.; Ciancanelli, R.; Noll, B. C.; Curtis, C. J.; Dubois, D. L.; Dubois, M. R. Organometallics 2002, 21, 4833−4839. (b) Miller, A. J. M.; Labinger, J. A.; Bercaw, J. E. Organometallics 2011, 30, 4308− 4314. (45) DuBois, D. L.; Berning, D. E. Appl. Organomet. Chem. 2000, 14, 860−862. (46) (a) Matsubara, Y.; Fujita, E.; Doherty, M. D.; Muckerman, J. T.; Creutz, C. J. Am. Chem. Soc. 2012, 134, 15743−15757. (b) Ellis, W. W.; Miedaner, A.; Curtis, C. J.; Gibson, D. H.; DuBois, D. L. J. Am. Chem. Soc. 2002, 124, 1926−1932. (c) Konno, H.; Kobayashi, A.; Sakamoto, K.; Fagalde, F.; Katz, N.; Saitoh, H.; Ishitani, O. Inorg. Chim. Acta 2000, 299, 155−163. (47) In the presence of water, the pKa of weak acids are expected to decrease: (a) Kütt, A.; Leito, I.; Kaljurand, I.; Soovali, L.; Vlasov, V. M.; Yagupolskii, L. M.; Koppel, I. A. J. Org. Chem. 2006, 71, 2829− 2838. (b) Kaupmees, K.; Kaljurand, I.; Leito, I. J. J. Phys. Chem. A 2010, 114, 11788−11793. (c) McCarthy, B. D.; Martin, D. J.; Rountree, E. S.; Ullman, A. C.; Dempsey, J. L. Inorg. Chem. 2014, 53, 8350−8361. (48) (a) Tilset, M.; Parker, V. D. J. Am. Chem. Soc. 1989, 111, 6711− 6717. (b) Wang, D.; Angelici, R. J. J. Am. Chem. Soc. 1996, 118, 935− 942. (c) Ryan, O. B.; Tilset, M.; Parker, V. D. J. Am. Chem. Soc. 1990, 112, 2618−2626. (49) Wayner, D. D. M.; Parker, V. D. Acc. Chem. Res. 1993, 26, 287− 294. (50) (a) Kelly, C. A.; Rosseinsky, D. R. Phys. Chem. Chem. Phys. 2001, 3, 2086−2090. (b) Pearson, R. G. J. Am. Chem. Soc. 1986, 108, 6109−6114. (51) (a) Aqueous-Phase Organometallic Catalysis: Concepts and Applications; Cornils, B., Herrmann, W. A., Eds.; Second Edition; Wiley−VCH: Weinheim, Germany, 2005. (b) Yoshida, T.; Kamato, K.; Tsukamoto, M.; Iida, T.; Schlettwein, D.; Wöhrle, D.; Kaneko, M. J. Electroanal. Chem. 1995, 385, 209−225. (c) Inglis, J. L.; MacLean, B. J.; Pryce, M. T.; Vos, J. G. Coord. Chem. Rev. 2012, 256, 2571−2600. (d) Savéant, J.-M. Chem. Rev. 2008, 108, 2348−2378. (52) (a) Fong, H.; Peters, J. C. Inorg. Chem. 2015, 54, 5124−5135. (b) Estes, D. P.; Vannucci, A. K.; Hall, A. R.; Lichtenberger, D. L.; Norton, J. R. Organometallics 2011, 30, 3444−3447. (53) Noviandri, I.; Brown, K. N.; Fleming, D. S.; Gulyas, P. T.; Lay, P. A.; Masters, A. F.; Phillips, L. J. Phys. Chem. B 1999, 103, 6713− 6722. (54) Krejcik, M.; Danek, M.; Hartl, F. J. Electroanal. Chem. Interfacial Electrochem. 1991, 317, 179−187.

(55) (a) SMART Software Users Guide, Version 5.1; Bruker Analytical X-ray Systems, Inc.: Madison, WI, 1999. (b) SAINT Software Users Guide, Version 7.0; Bruker Analytical X-ray Systems, Inc.: Madison, WI, 1999. (c) Sheldrick, G. M. SADABS, Version 2.3; Bruker Analytical Xray Systems, Inc.: Madison, WI, 2000. (d) Sheldrick, G. M. SHELXTL, Version 8.32; Bruker Analytical X-ray Systems, Inc.: Madison, WI, 2014. (e) International Tables for X-ray Crystallography, Vol. C; Kluwer Academic Publishers: Dordrecht, The Netherlands, 1992.

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