Article Cite This: Organometallics XXXX, XXX, XXX−XXX
Considering a Possible Role for [H-Fe4N(CO)12]2− in Selective Electrocatalytic CO2 Reduction to Formate by [Fe4N(CO)12]− Atefeh Taheri,† Natalia D. Loewen,† David B. Cluff, and Louise A. Berben* Department of Chemistry, University of California Davis, Davis, California 95616, United States S Supporting Information *
ABSTRACT: [Fe4N(CO)12]− is a first-row transition element electrocatalyst that selectively produces C−H bonds to give formate from CO2 in water at −1.2 V vs SCE. We present a thermochemical analysis which probes the possibility that [HFe4N(CO)12]2− ((H-1)2−) is an intermediate in this process: we show that (H-1)2− is accessible at −1.2 V vs SCE, but if it were formed, we predict that it would generate H2. [Fe4N(CO)12]3− and (H-1)2− were interrogated spectroscopically, and the product of CO loss, [Fe4N(CO)9(μ-CO)2]3−, was synthesized and characterized. Ultimately, we demonstrate that (H-1)2− is an unlikely participant in the catalytic transformation of CO2 to formate.
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INTRODUCTION We have previously reported that [Fe4N(CO)12]− (1−) is an electrocatalyst for reduction of CO2 into formate under mild conditions of pH 7 buffered aqueous solution and an applied potential of just −1.2 V vs SCE.1 The catalyst is stable for over 24 h in controlled-potential electrolysis experiments, and this provides a study system for selective conversion of CO2 into formate under aqueous conditions or in MeCN solution. One other first-row electrocatalyst selective for formate production has been reported by Artero and co-workers, and that example is a Co-P2N2 complex which operates at −2 V vs SCE in DMF solution over 1 h.2 A formate-selective PCP-Ir catalyst has also been reported in water.3 These catalysts offer a chance to investigate mechanisms of C−H bond formation with CO2 and to elucidate features needed to move forward with nextgeneration electrocatalysts for accessing hydrogenated CO2derived products as solar fuels.4 Our previous work on 1− provided a glimpse of the reaction mechanism with reaction rate, order of reaction, and determination that a cluster hydride, [H-Fe4N(CO)12]− (H1)−, is a likely intermediate: (H-1)− can be formed by reduction and subsequent protonation of 1− (Scheme 1). In an ECCE-type mechanism, reaction of (H-1)− with CO2 is proposed to afford HCOO− and 1, which is quickly reduced back to 1− under the reaction conditions of −1.2 V vs SCE (Scheme 1, left). We acknowledged in that prior work that other reaction mechanisms are possible, and this paper investigates an ECEC mechanism in which further reduced hydride, [H-Fe4N(CO)12]2− ((H-1)2−), might be involved in formate production. In this alternative mechanism, the initial reduction and protonation of 1− occur to produce (H-1)−, but then mechanistic pathways diverge if instead reduction affords (H-1)2−, which can potentially react with CO2 to give formate © XXXX American Chemical Society
Scheme 1. Proposed Mechanisms for Reduction of CO2 to Formate by 1−: (Left) Possible ECCE Mechanism; (Right) Possible ECEC Mechanism
and 1− (Scheme 1, right). We ultimately determine that the ECCE mechanism proposed previously is more reasonable.
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RESULTS AND DISCUSSION We first synthesized [H-Fe4N(CO)12] ((H-1)) to determine whether the proposed intermediate (H-1)2− is accessible under the catalytic conditions for formate production of −1.2 V.5 In hexane or toluene, IR spectra show pure (H-1). In solvents with higher dielectric constant, (H-1) dissociates partially into 1− and protons, and IR spectra show both (H-1) and 1− (Figure S1). The cyclic voltammogram (CV) of the (H-1)/1− mixture was acquired in 0.1 M Bu4NPF6 MeCN solution and showed four reversible reduction events which correspond to Received: November 14, 2017
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DOI: 10.1021/acs.organomet.7b00824 Organometallics XXXX, XXX, XXX−XXX
Article
Organometallics the four redox couples: (H-1)0/−, (H-1)−/2−, (1)−/2−, and (1)2−/3−, at −0.57, −0.91, −1.23, and −1.6 V vs SCE (Figure 1). From these data it is apparent that under the reaction
Figure 2. Solid-state structure of Na3[Fe4N(CO)11] in Na3-2. Green, gray, blue, red, and orange ellipsoids represent Fe, C, N, O, and Na atoms, respectively. The solvent molecules coordinated to Na atoms are omitted for clarity.
would be 2.45 Å.9 We have not been able to obtain single crystals of 23− in the absence of sodium cations, and we believe the Lewis acidic cations are required to stabilize the bridging carbonyls and the electron-rich cluster. We next set out to estimate the hydricity value for (H-1)2−, because this value can be used as a predictor of whether (H1)2− will react with H+ to produce H2 or whether it could selectively generate C−H bonds from CO2. Hydricity is defined as the free energy for loss of H− from a metal complex, for example (eq 1):11
Figure 1. CVs of 0.1 mM 1− (black) and H-1 (red) in 0.1 M Bu4NPF6 MeCN solution (GC working electrode, 100 mV s−1).
conditions for electrocatalytic reduction of CO2 into formate, −1.2 V vs SCE, (H-1)2− can potentially be generated from (H1)−. A synthetic route to (H-1)2− is protonation of 13−; therefore, we attempted to synthesize 13− to acquire its spectroscopic signals for use in infrared spectroelectrochemical IR-SEC experiments where we generate (H-1)2− in situ. Addition of 2.4−5 equiv of sodium anthracenide to a green-brown THF solution of 1− afforded a dark red-brown solution of [(diglyme)2(THF)4.5)Na3][Fe4(N)(CO)9(μ-CO)2] (Na3-2), which contains [Fe4(N)(CO)9(μ-CO)2]3− (23−). The infrared spectrum of this solution includes bands at 1790 (s), 1774 (s), and 1756 (m) cm−1 (Figure S2). The bands at lower energy suggest that bridging CO ligands are present. A diffusion of pentane into a concentrated solution containing Na3-2 in THF was stored at −25 °C for 2 weeks and yielded black crystals suitable for single-crystal X-ray diffraction (Tables S1 and S2). The solid-state structure of Na3-2 contains the cluster [Fe4(N)(CO)11]3− (23−), which is distorted relative to 1− by the nitride moving out of the butterfly core (Chart 1). A similar
[HFe4N(CO)12 ]2 − ⇌ [Fe4N(CO)12 ]− + H− (H‐1)2 −
1−
ΔG°H− (1)
We have previously shown in an iron carbonyl cluster series of electrocatalysts that, in water, clusters with 15 < ΔG°H− < 24 kcal mol−1 are most selective for C−H bond formation with CO2. In MeCN the target range is 44 < ΔG°H− < 49 kcalmol−1.12 ΔG°H− = 24 and 44 kcal mol−1 are the hydricities of formate in water13 and MeCN,14 respectively. Outside of these ranges H2 evolution dominates either because it is kinetically faster (at low ΔG°H−) or because the metal hydride is not a strong enough hydride donor to react with CO2 (high ΔG°H−). Others have made experimental observations, and theoretical predictions, that support these proposed ranges.15 In aqueous solution buffered between pH 6.5 and pH 13, we have previously shown that (H-1)− can only be stabilized at pH 13. At lower pH values reaction with available protons produces H2 immediately. Further reduced (H-1)2− cannot even be generated at pH 13. As we previously demonstrated, these experimental observations set an upper limit on the hydricity of (H-1)2− ,and so we know that the ΔG°H− value of (H-1)2− is 44: (a) Izutsu, K. Acid-Base Dissociation Constants in Dipolar Aprotic Solvents; Blackwell Science: Oxford, U.K., 1990; IUPAC Chemical Data Series. (b) Schwesinger, R.; Schlemper, H. Angew. Chem., Int. Ed. Engl. 1987, 26, 1167−1169. (19) Tilset, M.; Parker, V. D. J. Am. Chem. Soc. 1989, 111, 6711− 6717. (20) Barton, B. E.; Rauchfuss, T. B. Inorg. Chem. 2008, 47, 2261− 2263. (21) Curtis, C. J.; Miedaner, A.; Ellis, W. W.; Dubois, D. L. J. Am. Chem. Soc. 2002, 124, 1918−1925. (22) Wiedner, E. S.; Chambers, M. B.; Pitman, C. L.; Bullock, R. M.; Miller, A. J. M.; Appel, A. M. Chem. Rev. 2016, 116, 8655−8692. (23) (a) Kolis, W.; Drezdzon, M. A.; Shriver, D. F. Inorg. Synth. 1989, 26, 246. (b) Strong, H.; Krusic, P. J.; Fillipo, J. S. Inorg. Synth. 1990, 28, 203. (24) 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. (25) (a) 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. (b) Bard, A. J.; Faulkner, L. R.; Swain, E.; Robey, C. Electrochemical Methods: Fundamentals and Applications, 2nd ed.; Wiley: New York, 2001. (26) Krejcik, M.; Danek, M.; Hartl, F. J. J. Electroanal. Chem. Interfacial Electrochem. 1991, 317, 179−187. (27) (a) SMART Software Users Guide, Version 5.1, Bruker Analytical X-Ray Systems: Madison, WI, 1999. (b) SAINT Software Users Guide, Version 7.0; Bruker Analytical X-Ray Systems: Madison, WI, 1999. (c) Sheldrick, G. M. SADABS, Version 2.03; Bruker Analytical X-Ray Systems, Madison, WI, 2000. (d) Sheldrick, G. M. SHELXTL Version 6.12; Bruker Analytical X-Ray Systems, Madison, WI, 1999.
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DOI: 10.1021/acs.organomet.7b00824 Organometallics XXXX, XXX, XXX−XXX
