Formation of Copper Catalysts for CO2 Reduction with High Ethylene

Apr 4, 2016 - SUNCAT Center for Interface Science and Catalysis, Department of ... Department of Physics, AlbaNova University Center, Stockholm ...
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Formation of Copper Catalysts for CO2 Reduction with High Ethylene/Methane Product Ratio Investigated with In Situ X‑ray Absorption Spectroscopy André Eilert,†,‡,∥ F. Sloan Roberts,†,‡,∥ Daniel Friebel,†,‡ and Anders Nilsson*,†,‡,§ †

SLAC National Accelerator Laboratory, 2575 Sand Hill Road, Menlo Park, California 94025, United States SUNCAT Center for Interface Science and Catalysis, Department of Chemical Engineering, Stanford University, 443 Via Ortega, Stanford, California 95305, United States § Department of Physics, AlbaNova University Center, Stockholm University, Roslagstullsbacken 21, S-10691 Stockholm, Sweden ‡

S Supporting Information *

ABSTRACT: Nanostructured copper cathodes are among the most efficient and selective catalysts to date for making multicarbon products from the electrochemical carbon dioxide reduction reaction (CO2RR). We report an in situ X-ray absorption spectroscopy investigation of the formation of a copper nanocube CO2RR catalyst with high activity that highly favors ethylene over methane production. The results show that the precursor for the copper nanocube formation is copper(I)-oxide, not copper(I)-chloride as previously assumed. A second route to an electrochemically similar material via a copper(II)−carbonate/hydroxide is also reported. This study highlights the importance of using oxidized copper precursors for constructing selective CO2 reduction catalysts and shows the precursor oxidation state does not affect the electrocatalyst selectivity toward ethylene formation.

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In general, oxidizing copper in chloride aqueous solution can lead to a broad variety of chemical compounds such as Cu2(OH)3Cl and CuCO3·Cu(OH)2, depending on the actual conditions such as ion concentrations and pH.19,20 It is thus difficult to predict what copper species is formed from literature data. Therefore, we investigate the formation mechanism of CuCubes by in situ Cu K-edge X-ray absorption spectroscopy (XAS). It is important that the characterization is carried out in situ to avoid unwanted chemical transformations, e.g., due to exposure to air. Additionally, we introduce a novel copper electrocatalyst with different precursor oxidation state. The formation of CuCubes is achieved by subjecting polycrystalline copper to potentiodynamic oxidation−reduction cycles (Figure 1 inset) in 0.1 M bicarbonate solution and 4 mM KCl. With the addition of KCl, the cyclic voltammograms (CVs) show two distinct features: an increased anodic current above 0.56 V and a cathodic wave at −0.3 V (see Roberts et al. for more details).16 It seems reasonable that CuCl might grow at oxidative potentials in a cubic structure, which subsequently transforms to Cu2O and is then reduced to Cu in the cathodic wave.16 However, Cu K-edge XAS spectra show the presence of Cu2O but no evidence for CuCl. Figure 1 shows spectra at open circuit potential (ocp) after different electrochemical treatments, as well as spectra of reference compounds. The spectrum of the reduced sample (Red.) is, as expected, the same as for copper metal.21 Furthermore, no changes between

he electroreduction of carbon dioxide (CO2RR) to useful chemicals can be used to solve energy storage problems inherent to renewable electricity production and is a promising path toward a sustainable fuel cycle with net-neutral CO2 emissions.1,2 Energy efficiency and product selectivity of the CO2RR is strongly dependent on the electrocatalyst material.3 In particular, copper has attracted a great deal of attention because of its unique ability to reduce carbon dioxide to multicarbon products, such as ethylene, in high yields.4,5 Furthermore, nanostructures with copper oxide precursors have proven to be very efficient, with high yields of alcohols and other multicarbon products.6−15 Previously we reported the details of a cubic copper electrocatalyst (“CuCubes”) with exceptionally high ethylene over methane selectivity at low overpotentials.16,17 For CO2 reduction, the onset potential for ethylene production is −0.6 V versus the reversible hydrogen electrode (RHE), and no methane is detected over the entire investigated potential range to −1.15 V. Because this catalyst is formed by adding chloride to the electrolyte and oxidizing, it was hypothesized that its precursor is probably cubic CuCl which subsequently converted to Cu2O. This was in accordance with Chen et al., who reported an electrocatalyst with similar morphology and selectivity derived from a CuCl precursor.18 This is of particular importance because these materials are the only nanostructured copper catalysts with high multicarbon selectivities not thought to be directly derived from an oxide, whereas other nonoxide derived nanoparticles proved to be comparatively much less selective.7 © XXXX American Chemical Society

Received: February 17, 2016 Accepted: April 4, 2016

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The Journal of Physical Chemistry Letters

potential was held constant over the duration of a spectral series. Figure 2 shows these time-dependent transformations for four different conditions. In Figure 2a, the potential was ramped from ocp to the onset of the oxidation reaction at 0.56 V. As the spectra indicate and as quantitatively addressed by a linear combination analysis, the sample transforms from metallic copper to Cu2O within 20 min. There is no indication of any intermediate CuCl formation on the time scale of the measurements. As shown on the right-hand side of Figure 2, the transferred electrochemical charge and the oxidation state correlate well with each other, showing an oxidation of the entire film. A small, further increase of the potential to 0.6 and 0.7 V did not change the observed oxidation state of the residual film. The excess charge that is passed while holding at oxidizing potentials, without changing oxidation state, is suggestive of further copper oxidation to a species that dissolved into the electrolyte, which is in accordance with a decreasing edge jump (Figure S4). In Figure 2b, the potential was ramped from ocp to 0.7 V. Under this condition, the transformation to Cu2O, which took 20 min at 0.56 V, was completed within the first spectrum, i.e., within less than 5 min. After 10 min, the film started transforming into a Cu(II) species with reproducible spectra, similar to CuCO3, Cu(OH)2, and a [Cu(H2O)6]-salt (Figure S5). Thus, it is likely that the formed species contains both carbonate and hydroxide, such as the minerals malachite and azurite, which are the most stable compounds for similar electrolyte compositions according to Pourbaix.25 This Cu(II) species is usually not formed when making CuCubes because

Figure 1. XAS spectra at open circuit potential (ocp) of reduced CuCubes (Red.), CuCubes after the first, second, and third oxidation cycle at 0.7 V vs RHE (Ox.), and reference spectra (ref.). Cu reference from http://cars.uchicago.edu/~newville/ModelLib/;21 references for Cu2O, CuO, and CuCl from Ferrandon et al.23 Inset: Cyclic voltammogram from the second oxidation cycle and without KCl for comparison.

the initially reduced sample and the sample after several oxidation−reduction cycles were detected (Figure S1). In fact, peak positions stayed constant within the experimental resolution over the whole spectral region up to 9100 eV. A strain analysis using the “bond length with a ruler” concept22 resulted in no significant lattice strain; bond length changes, if any, are below the extended X-ray absorption fine structure detection limit of 0.01 Å (Figure S2). Furthermore, no residual oxide can be found in CuCubes. The spectrum of the oxidized sample (Ox. 1) can be well described by a linear combination of Cu and Cu2O spectra, where Cu2O is the dominant component. This behavior is reproducible over three oxidation−reduction cycles (Ox. 1, 2, and 3) that are used in a standard procedure for making CuCubes.16 A linear combination analysis (LCA) gives fractions of 85, 71, and 99% Cu2O for the first, second, and third oxidation cycle, respectively, which is in accordance with the fact that the second oxidation cycle reversed at a 37 mV lower voltage compared to the first and the third cycle. A comparison with reference spectra of CuCl and CuO,23 especially in the region between 8985 and 8990 eV, shows no indication of those species. It was not possible to obtain a linear combination fit with any physically meaningful concentration of CuCl; instead, including a CuCl component resulted in small negative CuCl concentrations in the oxidized samples. Cl K-edge XAS showed no change after oxidation (Figure S3 left), directly supporting no significant amounts of chlorine-containing species at the working electrode. A detection limit corresponding to a 0.3 nm thick CuCl film can be estimated from the Cl K-edge difference spectrum (Figure S3 right). No indication was found for any Cu(II) species, which exhibit no peak below 8985.0 eV.23,24 Cycling the potential to make CuCubes and measuring at open circuit potential does not investigate their formation mechanism under operating conditions. Therefore, to directly measure spectra while the sample is forming CuCubes, the

Figure 2. Evolution of XAS spectra (left), results of a two-component linear combination analysis (right, circles), and transferred electrochemical charge (right, solid line). (a) Slow oxidation in the presence of Cl; I, potential E ramped from ocp to 0.56 V; II, E = 0.56 V; III, E = 0.6 V; IV, E = 0.7 V. (b) Oxidation in the presence of Cl at 0.7 V. (c) Oxidation in the absence of Cl at 0.7 V. (d) Reduction in the presence of Cl at 0.35 V. All potentials vs RHE. 1467

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The Journal of Physical Chemistry Letters the potential is not held oxidative, but is instead cycled, leading to Cu2O as the precursor, as indicated in Figure 1. The oxidation of the film was also investigated in the absence of KCl (Figure 2c). After an anodic ramp and holding at 0.7 V, no oxidation was observable within the first 15 min. After that time, the film was directly transformed to the same Cu(II)carbonate/hydroxide species as in Figure 2b within the next 15 min. The existence of a waiting time before the oxidation starts was reproducible, with waiting times between 10 and 20 min observed in four repetitions. This points toward the growth of a thin passivation layer, which finally breaks and allows rapid oxidation as observed in pitting corrosion. Spectral concentrations of Cu2O of up to 2% could be added into the LCA of the spectra before rapid oxidation started without significant disagreement. Furthermore, a decreasing edge jump during rapid oxidation points toward additional copper dissolution (Figure S4). The direct transformation to Cu(II), i.e., bypassing the Cu(I) species observed when oxidizing with KCl, was also reproducible for pristine Cu samples. However, Cu2O concentrations of up to 20% were observed in samples that had already undergone previous oxidation−reduction cycles. Interestingly, a slow oxidation at 0.7 V turned out to be crucial. Holding the potential at 0.9 V instead led to copper stripping and no formation of any stable, solid oxidized copper compound. Figure 2d shows the spectral evolution during the slow reduction of the Cu2O precursor over 20 min at the onset potential for reduction at 0.35 V. The spectra change almost linearly back to metallic copper, and the final spectrum appears to be the same within experimental resolution as the pristine sample and the reduced CuCubes in Figure 1. This observation also holds true for the Cu(II)-carbonate/hydroxide precursor, reducing the Cu(II) species back to metal leads to the same spectrum as the pristine sample and the reduced CuCubes (Figure S6). This indicates no residual oxide or lattice strain in the reduced catalyst, as is further depicted in Figure S2. Online electrochemical mass spectrometry (OLEMS) during CO2RR using the Cu(II)-carbonate/hydroxide derived electrocatalyst (made by oxidation under potentiostatic conditions at 0.7 V without KCl) shows no significant differences to CuCubes (made by potentiodynamic oxidation−reduction cycles with KCl) (Figure 3a). The onset potential for CO2 reduction is at −0.6 V with an observed mass spectrometer signal ratio between ethylene and methane greater than 100:1. Note that OLEMS is not suitable for absolute, quantitative measurements because it records only gas-phase products and additional hydrogen is produced simultaneously. In contrast, untreated, polycrystalline copper shows a more negative onset potential for CO2 reduction at −0.8 V and no methane suppression. Both CuCubes and the Cu(II)-carbonate/ hydroxide derived Cu possess an involved nanostructure with high surface area, where the structures in Cu(II)-carbonate/ hydroxide derived Cu are generally larger than those in CuCubes (Figure 3b,c). There is an ongoing debate which structural properties make certain copper nanostructures superior in terms of electrochemical activity and selectivity toward multicarbon products compared to other, more inactive copper nanostructures. Hypotheses include the influence of sites with increased CO binding energy,26 grain boundaries,27 undercoordinated sites,28 residual oxides,11 and an increase of local pH.29,30 CuCubes were assumed to be unique because these were thought to be made from a CuCl precursor, whereas most other active

Figure 3. (a) OLEMS during CO2RR while cycling the potential using polycrystalline Cu, Cu2O derived CuCubes, and Cu(II)-carbonate/ hydroxide derived Cu; (b) scanning electron microscopy (SEM) images of CuCubes; and (c) the Cu(II)-carbonate/hydroxide derived electrocatalyst.

nanostructures were made from an oxide precursor. However, we were able to show that CuCubes are instead directly Cu2O derived and no solid CuCl is formed. Due to the exclusive formation of Cu(I)-oxide in the presence of KCl and Cu(II)-carbonate/hydroxide in its absence, it can be speculated that Cl− changes the oxidation mechanism, e.g., by the formation of transient Cu−Clcomplexes. These complexes could decrease the energy barrier to extract Cu atoms from their lattice, which could allow for Cu oxide growth via dissolution followed by precipitation, instead of a Cu/O place-exchange mechanism that could be predominant in the absence of Cl−. Such a different oxidation mechanism could also explain the increased anodic current in the presence of KCl (Figure 1 inset). Interestingly, the Cu(II)-carbonate/hydroxide derived Cu shows the same high electrochemical activity and selectivity toward ethylene formation during CO2RR as the Cu(I)-oxide derived CuCubes. This leads to the hypothesis that the actual morphology of the oxide and exposed facets of the reduced copper are of lower importance, as suggested by the large variety of different highly active nanostructures6−14,18,28,31 and addressed specifically by Baltrusaitis and co-workers.15 Because most of the previous studies employed Cu(I)-precursors, mostly Cu2O, our findings highlight the fact that the oxidation state of the precursor is insignificant for creating a copper electrocatalyst with increased activity and selectivity toward multicarbon products. Furthermore, an increase of local pH due to the consumption of protons at the cathode and mass-transfer limitations in nanostructures has shown to have a significant impact on CO2RR.29,30 Because the nanostructured samples show a higher current density compared to flat, polycrystalline copper (Figure S7), it can be assumed that an increase of local pH might indeed increase the yield of multicarbon products for the investigated nanostructured samples. However, there is a remarkable difference in current density between CuCubes and Cu(II)-derived Cu, which is not reflected in the OLEMS data (Figure 3a). The influence of local pH was studied more 1468

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The Journal of Physical Chemistry Letters thoroughly for CO reduction with CuCubes, where the solution pH can be varied.17 Because CuCubes showed exceptional yields for ethylene over methane even at pH 13 (where the local and solution pH are about the same), it was concluded that there is an important additional effect on-top of pH in CuCubes compared to other, flat Cu samples. Recently, the existence of residual oxides in similar samples under CO2RR conditions had been proposed,11 which were not detected in the thin XAS model samples investigated here. It should be noted that XAS is not a surface sensitive technique; thus, a residual oxide concentration in the order of a few nanometers would not have been detected. Furthermore, the XAS model samples were comparatively thin (200 nm or thinner after several oxidation−reduction cycles), which is necessary to prevent self-absorption in fluorescent mode XAS. Regarding the potentiodynamic growth of CuCubes on thick samples for OLEMS, a typical anodic charge of 25 mAs per oxidation wave was observed, which corresponds to a thickness of ca. 37 nm Cu being oxidized in the case of homogeneous oxidation. In the case of the potentiostatic growth of Cu(II)carbonate/hydroxide, thick samples showed a much higher transferred anodic charge corresponding to approximately 1000 nm oxidized Cu. Even though it is fairly certain that this thicker film also consists of Cu(II)-carbonate/hydroxide, it cannot be excluded that it has an enhanced stability of residual oxides. In conclusion, we investigated the formation mechanism of CuCubes, a highly efficient CO2RR catalyst for ethylene formation, by in situ Cu K-edge XAS. We were able to prove the existence of a Cu2O precursor, and no stable intermediate CuCl species were formed. This finding highlights the importance of oxygen-containing precursors to make copper electrocatalysts with high multicarbon selectivity and low overpotentials. By changing the formation parameters, we were able to introduce a novel Cu(II)-carbonate/hydroxide derived electrocatalyst with electrochemical properties similar to those of CuCubes. This points to the negligible importance of the precursor oxidation state. Finally, the active catalyst is metallic copper because no significant concentration of residual oxide was detected in the oxide-derived samples.

Radiation Lightsource (SSRL) on beamlines 7-3 and 14-3, respectively. XAS data analysis was performed using SixPack32 and algorithms written by the authors. Energy calibration was performed by comparing the first inflection point of the pristine, reduced sample spectrum to Ferrandon et al.23 Reference spectra from other sources were adjusted accordingly. The setup for online electrochemical mass spectrometry (OLEMS) is described in detail elsewhere.16 It consists of a mass spectrometer (SRS CIS 300) measuring electrochemical reaction products that enter the system through a porous Teflon frit (Porex 15−25 μm) placed near the surface of the copper working electrode. Reaction products were detected online while CVs were measured with the aforementioned parameters. The working electrodes were made of polycrystalline OHFC copper disks with 99.9% purity, with dimensions of 8 mm diameter and 2.5 mm height. The surfaces were characterized by scanning electron microscopy (FEI XL30 Sirion) in the Stanford Nanocharacterization Laboratory.

EXPERIMENTAL SECTION Working electrodes were made by the Microfabrication Shop in the Stanford Nano Shared Facilities at Stanford University. A 5 nm Ti adhesion layer and a 200 nm Cu film were deposited on 12 × 12 mm2, 250 μm thick glassy carbon windows (Structure Probe) using electron beam physical vapor deposition. For Cl K-edge XAS, ∼8 μm thick glassy carbon windows were used which were made by heating an 8 μm thick Kapton foil for 1 h at 1000 °C in a nitrogen atmosphere. The working electrode was glued with the Cu side facing inward on a hole in the wall of a 60 mL HDPE bottle (Figure S8) that was filled with an aqueous electrolyte solution of 4 mm KCl (>99.9995%, Fluka) and 0.1 m KHCO3 (>99.99%, Sigma-Aldrich) purged with CO2 (99.999%). Ultrapure water from a Millipore Gradient System with a resistivity >18 MΩ·cm was used. Cyclic voltammograms (CVs) were measured using a Ag/AgCl reference electrode (Innovative Instruments), a boron-doped diamond counter electrode (CCL Diamond), and a Biologic VSP 200 potentiostat scanning at a rate of 5 mV s−1. For Cl K-edge XAS, the HDPE laboratory bottle was kept in a helium atmosphere during the measurements. Cu and Cl K-edge XAS measurements were performed at the Stanford Synchrotron

ACKNOWLEDGMENTS This work was supported by the Air Force Office of Scientific Research through the MURI program under AFOSR Award No. FA9550-10-1-0572 and the Global Climate Energy Project at Stanford University. We thank Matthew Latimer and Erik Nelson for beamline support at SSRL beamlines 7-3 and 14-3. We thank Tom Carver from the Microfabrication Shop of the Stanford Nano Shared Facilities (SNSF) for preparing the thinfilm samples. SEM was performed in the Nanocharacterization Laboratory of SNSF. We thank Jürg Osterwalder, Filippo Cavalca, Anders F. Pedersen, and Hirohito Ogasawara for fruitful discussions.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpclett.6b00367. Sketch of the experimental setup, current density plot, lattice strain analysis, and more detailed XAS spectra (PDF)

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AUTHOR INFORMATION

Corresponding Author

*Stockholms universitet, Kemisk Fysik, 106 91 Stockholm, Sweden. E-mail: [email protected]. Author Contributions ∥

A.E. and F.S.R. contributed equally to this work.

Notes

The authors declare no competing financial interest.

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