# Patent

## Title
Catalysts for low temperature electrolytic CO2 reduction

## Abstract

A method for electrochemically reducing CO2 is provided. A cathode is provided, wherein the cathode comprises a conductive substrate with a catalyst of a metal and a metal oxide based coating on a side of the cathode. An anode is spaced apart from the cathode. An ionic transport is provided between the anode and cathode. The cathode is exposed to CO2 and H2O. The anode is exposed to H2O. A voltage is provided between the cathode and anode.

## Background
CROSS REFERENCE TO RELATED APPLICATIONS
This application claims priority under 35 U.S.C. §119 from U.S. Provisional Patent Application No. 61/511,824, filed Jul. 26, 2011, entitled CERIA-BASED ELECTROREDUCTION CATALYSTS FOR LOW-TEMPERATURE ELECTROLYTIC SYNGAS PRODUCTION and U.S. Provisional Patent Application No. 61/579,422, filed Dec. 22, 2011, entitled CATALYSTS FOR LOW TEMPERATURE ELECTROLYTIC CO2 REDUCTION, which are incorporated herein by reference for all purposes.


BACKGROUND OF THE INVENTION
This invention relates generally to the reduction of CO2. Sustainable production of C-based fuel requires using renewable energy to power the reductive fixation of CO2. Coupling renewable electricity to an electrolytic device is an attractive strategy for this goal because it enables the use of multiple renewable energy sources and independent optimization of catalysis. Solid oxide electrolytic cells reduce CO2 to CO efficiently at high current densities, but require operating temperatures of 750-900° C. and cannot access other products.
Materials that catalyze electrochemical CO2 reduction under mild conditions would enable the development of electrolyzers that operate at more convenient temperatures and provide access to alternative reduction products such as formic acid, alcohols and hydrocarbons. Researchers over the past three decades have identified several materials that are capable of reducing CO2 electrochemically in aqueous solutions, but none that is efficient and stable enough for practical use. In general, available electrodes suffer from one or more of three major problems: 1) a requirement for excessive reducing potentials (“overpotentials”) to reduce CO2 in preference to reducing H2O, resulting in low energetic efficiency; 2) rapid loss of CO2 reduction activity resulting from electrode poisoning; 3) production of multiple CO2 reduction products with little selectivity. There is a pressing need to discover and develop new electrochemical CO2 reduction catalysts in order for sustainable fuels to be a significant contributor to a renewable energy economy.
SUMMARY OF THE INVENTION
In accordance with the invention, a method for electrochemically reducing CO2 is provided. A cathode is provided, wherein the cathode comprises a conductive substrate with a catalyst of a metal and a metal oxide based coating on a side of the cathode. An anode is spaced apart from the cathode. An ionic transport is provided between the anode and cathode. The cathode is exposed to CO2 and H2O. The anode is exposed to H2O. A voltage is provided between the cathode and anode.
In another manifestation of the invention, a method for electrochemically reducing CO2 is provided. A coating is formed on a cathode by heating a metal layer of the cathode in air, electrochemically oxidizing the metal layer of the cathode, or by a metal oxide deposition to form a metal and metal oxide interface. An anode is spaced apart from the cathode. An ionic transport is provided between the anode and cathode. The coating is exposed to CO2 and H2O. The anode is exposed to H2O. A voltage is provided between the cathode and anode.
In another manifestation of the invention an apparatus, for electrochemically reducing CO2 is provided. An anode is provided. An oxidized cathode is spaced apart from the anode. A chamber for exposing the anode and oxidized cathode to at least one electrolyte is adjacent to the anode and oxidized cathode. A gas chamber for exposing the oxidized cathode to CO2 is adjacent to the oxidized cathode. A CO2 source for providing CO2 to the gas chamber is connected to the gas chamber.
The invention and objects and features thereof will be more readily apparent from the following detailed description and appended claims when taken with the drawings.



BRIEF DESCRIPTION OF THE DRAWINGS
FIG. 1A shows the XPS spectra of untreated Sn foil before and after electrolysis and Sn foil after etching in HBr.
FIG. 1B is a plot of total current density vs time, CO faradaic efficiency vs time, and overall HCO2H faradaic efficiency at −0.7 V vs RHE in CO2-saturated 0.5 M NaHCO3 for unetched Sn.
FIG. 1C is a plot of total current density vs time, CO faradaic efficiency vs time, and overall HCO2H faradaic efficiency for untreated Sn at −0.7 V vs RHE in CO2-saturated 0.5 M NaHCO3 for etched Sn.
FIG. 2A depicts the bulk electrolysis trace at −0.7 V in NaHCO3/CO2 electrolyte for a Ti cathode before and after the addition of 1 mM SnCl2 to the electrolyte.
FIG. 2B shows SEM images of a Ti electrode before and after deposition showing the formation of a porous, particulate film with  ̃100 nm-diameter pieces atop a more uniform layer.
FIG. 2C is a high resolution Sn 3d5/2 XPS of a Sn/SnOx catalyst removed 30 mM or 12 h after the addition of Sn2+.
FIG. 2D provides graphs of XRD patterns showing Sn0, SnO2, and Ti peaks after 30 min or 12 h.
FIGS. 3A-C shows the comparison of CO2 reduction catalysis for unexcited Sn foil and in situ deposited Sn/SnOx thin film electrodes.
FIGS. 4A-E shows the total geometric current density (jtot) vs time, the faradaic efficiency (FE) for CO vs time and the overall FE for HCO2H for the polycrystalline Cu electrode and several of the annealed electrodes with progressively thicker initial Cu2O layers at −0.5 V vs the reversible hydrogen electrode.
FIG. 4F shows the average FE for CO vs the amount of charge required to reduce the Cu2O layer per electrode area.
FIGS. 5A-F show the scanning electron microscopy (SEM) images, X-ray diffraction (XRD) patterns, and high-resolution Cu 2p X-ray photoelectron spectroscopy (XPS) spectra for a Cu electrode after annealing procedure and after subsequent CO2 reduction electrolysis.
FIGS. 6A-C show the total current densities and faradaic efficiencies for the major products for a Cu electrode annealed at 500° C. for 12 h and for polycrystalline Cu.
FIG. 7 shows Tafel data for a Cu electrode annealed at 500° C. for 12 h and Tafel data for polycrystalline Cu.
FIG. 8 is a high level flow chart of an embodiment of the invention.
FIGS. 9A-B are enlarged cross-sectional views of part of a conductive substrate with a metal coating, forming part of a cathode.
FIG. 10 is a schematic view of an electrolyzer that may be used in an embodiment of the invention.



DETAILED DESCRIPTION OF ILLUSTRATED EMBODIMENTS
Tin
Metal electrodes have been the focus of extensive CO2 electroreduction studies in aqueous solutions at ambient temperature. Sn has attracted considerable interest because it is one of the most active metals and its low cost is amenable to large-scale use. Despite its appeal relative to other electrodes, the energy efficiency of Sn is too low for practical electrolysis. Sn is reported to require at least 0.86 V of overpotential to attain a CO2 reduction partial current density of 4-5 mA/cm2 in an aqueous solution saturated with 1 atm of CO2. It is generally assumed that the bare Sn surface is the catalytically active surface for CO2 reduction. The large overpotential required for CO2 reduction is thought to result from the barrier associated with the initial e− transfer to form a CO2•− intermediate that is poorly stabilized by the Sn surface. This mechanistic scenario is commonly invoked for many metal electrodes.
In an embodiment of the invention, SnOx is essential to CO2 reduction catalysis on Sn. This may be shown by demonstrating that removal of SnOx from a Sn electrode results in nearly exclusive H2 evolution activity. This insight is subsequently applied to prepare a composite Sn/SnOx thin film catalyst that exhibits greatly enhanced CO2 reduction activity relative to a typical Sn electrode.
To evaluate the importance of SnOx on the surface of Sn in CO2 reduction, we compared the activity of Sn electrodes that had been etched in strong acid to the activity of untreated electrodes. In both cases, new pieces of high purity Sn foil (99.998%) were used. The surface of the untreated foil was examined by XPS to characterize the native SnOx layer. FIG. 1A shows the XPS spectra of untreated Sn foil before and after electrolysis (left) and Sn foil after etching in HBr (right). The curves are combinations of two Gaussian/Lorentzian curves at 486.5 eV and 484.7 eV. FIG. 1B is a plot of total current density vs time (indicated by the line), CO faradaic efficiency vs time (indicated by the ▪ points) and overall HCO2H faradaic efficiency for untreated Sn at −0.7 V vs RHE in CO2-saturated 0.5 M NaHCO3 for unetched Sn. FIG. 1C is a plot of total current density vs time (indicated by the line), CO faradaic efficiency vs time (indicated by the ▪ points) and overall HCO2H faradaic efficiency for untreated Sn at −0.7 V vs RHE in CO2-saturated 0.5 M NaHCO3 for etched Sn.
The high resolution Sn 3d5/2 spectrum was fit to two peaks at 486.5 eV and 484.7 eV that correspond to Sn4+/2+ (SnOx) and Sn0, respectively. The ratio of corrected peak areas for SnOx to Sn0 is 95:5, indicating the presence of a >5 nm native SnOx layer.
Etched electrodes were prepared by immersing the Sn foil in 24% HBr at 90° C. for 10 min. An XPS spectrum of the etched electrode taken immediately after removal from the HBr solution exhibited a SnOx:Sn0 ratio of 17:83 (FIG. 1A). The residual oxide observed on this electrode is likely due to oxide regrowth in the brief exposure to air upon transferring to the XPS chamber, as assessed by independent XPS experiments with a sputtered electrode. For electrolysis experiments, etched electrodes were rinsed with deionized water at the conclusion of the etching procedure and used immediately to minimize oxide regrowth.
The electrolyses were performed in an H-cell in 0.5 M aqueous NaHCO3 saturated with CO2 (“NaHCO3/CO2”) at a potential of −0.7 V vs the reversible hydrogen electrode (RHE; all potentials are referenced to this electrode). The headspace of the cathodic compartment was continuously purged with CO2 into the sampling valve of a gas chromatograph (GC), enabling periodic quantification of the gas phase products. FIG. 1B shows the total geometric current density (jtot) vs time and the faradaic efficiency for CO production at various time points for an untreated Sn electrode. The electrode exhibits a current density of 0.4-0.6 mA/cm2 and a steady-state faradaic efficiency for CO of 5-10%. NMR analysis of the electrolyte at the conclusion of the experiment indicates 19% faradaic efficiency for HCO2H; the remainder of the current is accounted for by H2 formation. This CO2 reduction activity is consistent with the best reported activity for Sn at −1.06 V, taking into account the difference in overpotential. An electrode examined by XPS after a 12 h electrolysis at −0.7 V exhibited a SnOx:Sn0 ratio of 89:11, indicating that the native SnOx layer is stable to the reduction conditions (FIG. 1A).
Strikingly, an etched Sn electrode exhibits a much higher jtot of 3-4 mA/cm2, but very low faradaic efficiency for CO (0.5%) and HCO2H production (0.3%) (FIG. 1C). The higher jtot likely reflects a larger electrochemical surface area due to etching. Despite the higher surface area, the geometric partial current density for CO2 reduction is lower for the etched Sn electrode (24-32 μA/cm2) than the untreated Sn electrode (92-140 μA/cm2) due to the much lower faradaic efficiency. Very low (<1%) CO2 reduction faradaic efficiencies on etched Sn are also observed over a range of potentials from −0.5 V to −1.0 V. Thus, etched Sn is a moderately efficient H2 evolution catalyst, but is essentially inactive for CO2 electroreduction. Similar results were obtained if Sn electrodes were etched by polarizing at −3 V in HCl solution instead of treating with hot HBr solution.
Together, the XPS and electrolysis results indicate that removal of the native SnOx layer from a Sn electrode suppresses CO2 reduction activity such that H2 evolution accounts for >99% of the current density. The small residual CO2 reduction activity observed on etched Sn likely reflects the growth of a small amount of SnOx on the etched electrode before the start of electrolysis.
Based on these results, we hypothesized that the simultaneous deposition of Sn0 and SnOx on an electrode surface would result in a material with enhanced Sn—SnOx contact that is consequently a more active catalyst for CO2 reduction than a typical Sn foil electrode with a native SnOx layer. Accordingly, we sought electrodeposition conditions in which the hydrolysis of Sn2+ by cathodically generated OH− would take place concurrently with the reduction of Sn2+ to Sn0 (E0=−0.1375 V vs NHE). As described below, deposition on Ti electrodes under the same conditions used for CO2 electroreduction proved to be particularly effective.
FIG. 2A depicts the bulk electrolysis trace at −0.7 V in NaHCO3/CO2 electrolyte for a Ti cathode before and after the addition of 1 mM SnCl2 to the electrolyte. Prior to the addition of Sn2+, the Ti electrode exhibits a current density of  ̃10 μA/cm2 with very little detectable CO2 reduction. Addition of Sn2+ results in a sharp rise in the current density to a steady-state value of  ̃1.8 mA/cm2 and the formation of a grey deposit on the electrode surface. The current density is stable for >10 h and corresponds to >85% CO2 reduction with the remainder accounted for by H2 evolution. Nearly identical results are obtained if Sn(OTf)2 is used instead of SnCl2, indicating that Cl+ is not necessary for catalyst formation.
The composition and structure of the electrodeposited catalyst were characterized by a combination of scanning electron microscopy (SEM), XPS and powder x-ray diffraction (XRD). A catalyst was prepared via in situ deposition as described above and removed from the electrolyte 30 min after the addition of Sn2+. FIG. 2B shows SEM images of a Ti electrode before (left) and after (right) deposition showing the formation of a porous, particulate film with  ̃100 nm-diameter pieces atop a more uniform layer. FIG. 2C is a high resolution Sn 3d5/2 XPS of a Sn/SnOx catalyst removed 30 min (left) or 12 h (right) after the addition of Sn2+. XPS analysis indicates a SnOx:Sn0 ratio of 93:7, similar to the ratio observed for Sn foil electrodes with a native SnOx layer. FIG. 2D provides graphs of XRD patterns showing Sn0 (▪), SnO2 (★) and Ti (●) peaks after 30 min or 12 h. In the XRD pattern of this electrode, strong Sn0 peaks are observed along with small peaks that correspond to SnO2. The latter are absent for a Sn foil electrode with a native SnOx. For comparison, a separate catalyst film was prepared and removed for analysis 12 h after the addition of Sn2+. The XPS spectrum, shown in FIG. 2C, and XRD pattern for this electrode are very similar to those of the sample removed after 30 min. Together, these results indicate that a composite Sn/SnOx material is formed under the deposition conditions.
The electrodeposited catalyst (hereafter referred to as “Sn/SnOx”) exhibits greatly enhanced CO2 reduction catalysis compared to a typical Sn foil electrode with a native SnOx layer. For both electrodes, CO, HCO2H and H2 together account for >99% of the reduction products in NaHCO3/CO2 electrolyte. To compare the activities of Sn foil and Sn/SnOx, we measured their partial current densities for CO and HCO2H at selected potentials between −0.5 and −0.7 V. Comparison of CO2 reduction catalysis for Sn foil and in situ deposited Sn/SnO, thin film electrodes are illustrated in FIGS. 3A-C. FIG. 3A shows Tafel plots for HCO2H production. FIG. 3B shows Tafel plots for CO production. FIG. 3C is a bar graph showing Faradaic efficiencies for HCO2H and CO at various potentials. These data were obtained by performing stepped-potential electrolyses with periodic quantification of the gaseous products by GC and removal of aliquots after each step for NMR analysis.
For Sn foil, approximate Tafel slopes of 74 mV/dec and 72 mV/dec are observed for HCO2H and CO production, respectively. Similar Tafel slopes are observed for HCO2H (67 mV/dec) and CO (77 mV/dec) production on Sn/SnOx, however the geometric partial current densities are 7-8-fold higher than for Sn foil. The higher geometric current densities on Sn/SnOx are not simply the result of greater electroactive surface area, as indicated by cyclic voltammetry and the dramatic differences in faradaic efficiencies for Sn foil and Sn/SnOx. Over the range of potentials used for Tafel analysis, the CO faradaic efficiencies are 4-fold higher and the HCO2H faradaic efficiencies are 2-3-fold higher on Sn/SnOx than on untreated Sn foil.
The Tafel slopes for HCO2H and CO production on both Sn foil and Sn/SnOx are inconsistent with CO2 reduction mechanisms that proceed through an initial rate-determining 1 e− transfer to CO2. Such a mechanism would result in a 118 mV/dec slope. The observed slopes are instead much closer to 59 mV/dec, which supports mechanisms in which there is a reversible 1 e− transfer to CO2 to form CO2•− prior to a chemical rate-determining step. Possibilities for the chemical rate-determining step include protonation of CO2•− or migration to an alternative site on the electrode surface. Competing rate-determining steps, such as protonation at C vs O of CO2•−, may determine the HCO2H vs CO selectivity.
The Tafel data, combined with the absence of appreciable CO2 reduction activity on etched Sn, suggest that SnOx enables CO2 reduction to occur by stabilizing CO2•−. At present, we cannot determine whether reduction takes place at the interface between Sn0 and SnOx or on the SnOx surface directly. In the absence of SnOx to stabilize CO2•−, Sn0 only catalyzes H2 evolution because the 1 e− transfer to CO2 is prohibitively slow. The higher CO2 reduction partial current density and faradaic efficiency on Sn/SnOx relative to Sn foil with a native SnOx layer are therefore indicative of a greater density of active sites for CO2 reduction and a higher ratio of these sites to H2 evolution sites for the in situ deposited catalyst.
The CO2 reduction activity of Sn/SnOx, as indicated by the Tafel plots and faradaic efficiencies in FIGS. 3A-C, compares favorably to all polycrystalline metal electrodes in aqueous electrolytes with the exception of Au, which is comparably active initially, but subject to rapid deactivation. Improving CO2 and ion mass transport by incorporating Sn/SnOx in a flow cell and/or a gas diffusion electrode may enable increasing the current density by 1-2 orders of magnitude without large overpotential increases. Elucidating the detailed mechanistic role of SnOx in mediating electron transfer to CO2 is an important objective toward this goal. Moreover, the importance of SnOx to CO2 reduction on Sn surfaces raises the possibilities that metal oxides may be involved in CO2 reduction pathways on other metal electrodes and that the preparation of alternative metal/metal oxide composites may yield additional CO2 reduction catalysts with superior activity.
Copper
Polycrystalline Cu has been the focus of most CO2 reduction studies because it is one of the best available catalysts and is capable of producing hydrocarbon products. Although mechanistic studies have yielded valuable insights into the CO2 reduction pathways on Cu, the principal shortcomings of this electrode have not been addressed. Most significantly, the energetic efficiency of Cu is limited by the large overpotential (>0.7 V) required for CO2 reduction to outcompete H2O reduction. In addition, Cu electrodes rapidly lose their CO2 reduction activity unless stringently purified electrolytes are used, a requirement that is not compatible with scalable fuel synthesis.
Achieving efficient Cu-catalyzed CO2 reduction requires preparing Cu particles whose surfaces have active sites that are different from those on the surface of a polycrystalline Cu electrode. Electrochemical reduction of metal oxides provides one possible route to metal particles with altered surface structures. Researchers have previously used electrochemical methods including potential cycling and anodic pulses to form and subsequently reduce oxides on Cu electrodes. These treatments have resulted in increased hydrogen evolution activity in alkaline electrolytes and altered product selectivity at high overpotential in CO2 reduction electrolyses. While these studies provide evidence of altered electrocatalytic properties, substantial improvements to the energetic efficiency of CO2 reduction have not been observed. Researchers have also used copper oxide electrodes in CO2 reduction electrolyses. The oxides were reduced to Cu0 in situ during CO2 reduction catalysis, but only transient changes in the CO2 product distribution attributed to oxide catalysis were observed. Here we show that the CO2 reduction properties of Cu0 electrodes resulting from copper oxide reduction vary widely depending on the properties of the initial oxide layer. Reduction of thick Cu2O layers formed by high temperature annealing results in electrodes that catalyze energy-efficient CO2 reduction and are stable to the deactivation phenomena that plague bulk metal electrodes.
Electrodes were prepared by electropolishing pieces of polycrystalline Cu foil (99.9999%) in 85% phosphoric acid and subsequently annealing the electrodes in air at selected temperatures for variable amounts of time. The activities of these electrodes were compared to that of a polycrystalline Cu electrode in controlled potential electrolyses performed in CO2-saturated 0.5 M NaHCO3 electrolyte (“NaHCO3/CO2”) in a two-compartment electrolysis cell. The headspace of the cathodic chamber was continuously purged with CO2 into the sampling loop of a gas chromatograph (GC) to enable periodic quantification of the gas-phase products. The solution-phase products were quantified by NMR analysis of the electrolyte at the conclusion of the electrolyses.
FIGS. 4A-E shows the total geometric current density (jtot) vs time, the faradaic efficiency (FE) for CO vs time and the overall FE for HCO2H for the polycrystalline Cu electrode (FIG. 4A) and several of the annealed electrodes (FIG. 4B-E) with progressively thicker initial Cu2O layers at 0.5 V vs the reversible hydrogen electrode (RHE; all potentials are referenced to this electrode). The polycrystalline Cu electrode exhibited a jtot of  ̃100 μA/cm2, a FE for CO that declined from 10% at the start of the electrolysis to <2% over the course of 7 h and a FE for HCO2H of 3%. The majority of the current, >90%, was due to H2 evolution. These values are consistent with the previously measured activity for Cu in KHCO3 electrolytes. Annealing Cu at 130° C., the temperature used to prepare Cu2O electrodes for most previous studies, had very little effect on the activity under these conditions. The electrode annealed at 130° C. for 12 h (FIG. 4B) exhibited a jtot of 10 mA/cm2 during the first 4 s in which the thin Cu2O layer was reduced. Subsequently, the jtot and FEs were very similar to those of the polycrystalline electrode.
In contrast to these results, the electrodes annealed at higher temperatures exhibited larger jtot values and improved CO2 reduction FEs upon reduction of the Cu2O layer. The electrode annealed at 300° C. for 30 min exhibited an initial jtot of 10 mA/cm2 for 2 min as the Cu2O was reduced and subsequently a stable jtot of 1.0 mA/cm2. The FE for CO was 25% during the first hour of electrolysis before declining to 10% over 7 h; the FE for HCO2H on the reduced electrode was 5%. Further improvements were obtained by starting with a thicker Cu2O layer. After Cu2O reduction of the electrode annealed at 300° C. for 5 h, jtot reached a stable value of 1.3 mA/cm2, the FE for CO reached 35% and the FE for HCO2H was 24% (FIG. 4D). Annealing at 500° C. for 12 h resulted in an even thicker Cu2O layer and a stable jtot of 2.7 mA/cm2. This electrode produced CO with 40% FE and HCO2H with 33% FE. Notably, the FE for CO was maintained at 40% throughout the electrolysis, indicating not only efficient but also stable activity for CO2 reduction on this surface.
A plot of the average CO FEs for the annealed electrodes vs the amount of charge passed per electrode area (Q) in the Cu2O reduction is shown in FIG. 4F. The FEs increased with the amount of charge passed until reaching a plateau at 30-40% for Q≧ ̃5 C/cm2. Assuming bulk density of Cu2O on the electrode, 5 C/cm2 corresponds to a  ̃3 μm-thick layer. Together, these results demonstrate that a threshold thickness of the initial Cu2O layer is required to achieve both efficient and stable CO2 reduction catalysis for the electrode resulting from Cu2O reduction. Based on these results, electrodes prepared by annealing Cu at 500° C. for 12 h were selected for further characterization and CO2 reduction studies. FIGS. 5A-F show the scanning electron microscopy (SEM) images (FIGS. 5A,D), X-ray diffraction (XRD) patterns (FIGS. 5B,E), and high-resolution Cu 2p X-ray photoelectron spectroscopy (XPS) spectra (FIGS. 5C,F) for a Cu electrode after this annealing procedure (FIGS. 5A-C) and after subsequent CO2 reduction electrolysis (FIGS. 5D-F). After annealing, the SEM showed a dense array of rods with 100-1000 nm diameters on the electrode surface. These rods are the outermost portion of a thick Cu2O layer coating the electrode, as evidenced by the large Cu2O peaks and the near complete suppression of the Cu° peaks in the XRD pattern. The characteristic Cu2+ satellite peaks in the XPS spectrum are consistent with the presence of a thin (<10 nm) CuO layer coating the Cu2O. Following CO2 reduction electrolysis, SEM indicated that the rod morphology was intact, but smaller particles ( ̃20 nm) were embedded within the rods (FIG. 5D and FIG. S3). Only Cu0 peaks were observed in the XRD pattern, FIG. 5F. The Cu 2p XPS spectrum indicated the presence of Cu0 or Cu1+, but the peaks associated with Cu2+ in the spectra prior to electrolysis were absent. Together, these results indicate the complete reduction of the Cu2O layer, although we cannot rule out the presence of a thin, metastable Cu2O layer or other surface-bound Cu1+ species during electrocatalysis.
The electrochemically active surface area of a reduced electrode that had been annealed at 500° C. for 12 h was determined by measuring the double layer capacitance in 0.1 M HClO4 after CO2 reduction electrolysis. The capacitance was 13.9 mF/cm2, which is 475× larger than the capacitance of 29 μF/cm2 measured for a polycrystalline Cu electrode. This roughness factor is considerably larger than the difference in jtot between the two electrodes ( ̃30×), consistent with the difference in FEs between the two electrodes.
The presence of 100-1000 nm rods observed in FIG. 5D is not necessary for efficient CO2 reduction. Electrodes annealed at temperatures ≧500° C. for variable amounts of time exhibited very different morphological features on this length scale, but nonetheless comparable FEs for CO2 reduction at −0.5 V. These results suggest that the CO2 reduction efficiency of electrodes annealed at high temperatures is associated with a Cu particle surface or grain boundary structure that forms when suitably thick Cu2O layers are electrochemically reduced.
To further characterize the effect of high temperature annealing on the CO2 reduction activity of Cu, we measured the partial current densities for the reduction products at a variety of potentials between −0.2 V and −1.0 V in NaHCO3/CO2 using an electrode that had been annealed at 500° C. for 12 h (hereafter referred to as “annealed Cu”). The total current densities and faradaic efficiencies for the major products are shown in FIGS. 6A-C, which provides comparisons of electrocatalytic activities of polycrystalline Cu and Cu annealed at 500° C. for 12 h. FIG. 6A is a graph of total current density vs. potential. FIG. 6B is a graph of faradaic efficiencies for CO and HCO2H vs potential. FIG. 6C is a graph of faradaic efficiencies for CH4, C2H4 and C2H6 vs potential. Attempts to collect the corresponding data under identical conditions with polycrystalline Cu were unsuccessful due to the rapid degradation of catalytic activity. Instead, optimal data from previous studies with polycrystalline Cu at several potentials in 0.1 M KHCO3 are included for comparison.
The annealed Cu electrode exhibits a high efficiency for CO2 reduction at remarkably low overpotentials. A peak faradaic efficiency of 45% for CO production is obtained at potentials ranging from 0.3 V to 0.5 V, corresponding to 0.19 V to 0.39 V of overpotential for this product (FIG. 6B). By comparison, essentially no CO2 reduction to CO is observed for polycrystalline Cu in this potential range; the maximum efficiency for CO with polycrystalline Cu is 20%, which requires 0.8 V (η=0.69 V). Similarly, annealed Cu attains a peak faradaic efficiency for HCO2H production of 30% at potentials ranging from 0.45 V to 0.65 V (η=0.25 V to 0.45 V), whereas polycrystalline Cu requires −0.7 V to −0.9 V (η=0.5 V to 0.7 V) to attain a comparable faradaic efficiency (FIG. 6B).
At relatively negative potentials (<−0.6 V), annealed Cu catalyzes the reduction of CO2 to ethylene and ethane (FIG. 6C). In contrast, polycrystalline Cu produces only ethylene and methane at high overpotential. Previous work on Cu single crystals has shown that the ratio of ethylene to methane can be boosted by introducing (111) steps in the (100) basal plane, i.e. by using single crystal Cu electrodes with a high index face exposed to the solution. However, methane was never fully suppressed and no ethane was observed in these studies. These results indicate that the surface structures of the Cu particles produced by Cu2O reduction are distinct from the structures of the high index faces of Cu. We also note that no methanol was detected among the reduction products for annealed Cu at any potential examined here, in contrast to what has been reported for CO2 reduction catalysis with Cu electrodes annealed at lower temperatures.
The faradaic efficiencies for the hydrocarbon products on annealed Cu are low and H2 is the major product at high overpotentials. This difference relative to the lower overpotential regime most likely reflects the mass transport limitations at the high current densities observed (>10 mA/cm2) rather than the intrinsic selectivity of the electrode Improvements in mass transport by using a flow cell or gas diffusion electrode are expected to enable substantially higher CO2 reduction current densities without large overpotential increases.
To obtain insight into the mechanistic pathway(s) for CO2 reduction with annealed Cu, a plot of overpotential vs the log of the partial current density for CO production (a Tafel plot) was extracted from the data described above. The data are shown in FIG. 7 along with Tafel data for polycrystalline Cu. The plot for annealed Cu is linear over the range of overpotentials from 0.05 V to 0.3 V with a slope of 116 mV/decade. This slope is consistent with a rate-determining initial electron transfer to CO2 to form a surface-adsorbed CO2•− intermediate, a mechanism that is commonly invoked for metal electrodes. A similar slope is evident in the plot for polycrystalline Cu. The dramatic difference in FE between the two electrodes suggests that the Cu surfaces formed by reducing thick Cu2O layers enable formation of the CO2• intermediate while suppressing H2O reduction.
In summary, our results show that Cu particles prepared by reducing μm-thick Cu2O films catalyze the reduction of CO2 to CO and HCO2H with high faradaic efficiencies at exceptionally low overpotentials and produce C2 hydrocarbons to the exclusion of CH4 at high overpotentials. Electrodes with these characteristics can readily be prepared with high surface areas, enabling >1 mA/cm2 geometric current densities for CO2 reduction at <0.4 V overpotential and measurable CO2 reduction current densities at <0.1 V overpotential, levels of activity that were previously inaccessible with metal electrodes under comparable conditions. Furthermore, CO2 reduction with these electrodes is resistant to deactivation for at least several hours, a marked improvement over the rapid deactivation of polycrystalline Cu under identical conditions. We anticipate that elucidation of the surface structures of the Cu particles formed by reducing thick Cu2O layers will provide crucial insights into the structural requirements for preferential CO2 reduction and the formation of C2 products. In addition, this synthetic approach may prove useful for preparing additional electrocatalysts for CO2 reduction.
Embodiments of Implementation
To facilitate understanding of the invention, FIG. 8 is a high level flow chart of an embodiment of the invention. In this embodiment, a cathode with a catalyst metal is provided (step 804). A metal oxide coating is formed on the catalyst metal (step 808). The metal oxide coating and the catalyst metal form a metal and metal oxide coating, which may comprise a metal oxide coating over a metal coating or a single coating with both metal oxide particles and metal particles. An anode is spaced apart from the cathode (step 812). An ionic transport is provided between the anode and cathode (step 816). The cathode is exposed to CO2 and H2O (step 820). The anode is exposed to H2O (step 824). A voltage is provided between the cathode and anode (step 828). The voltage causes CO2 and H2O to be reduced to CO, H2, and O2. The CO and H2 may be converted to hydrocarbon or alcohol products.
In a specific embodiment of the invention, the cathode is formed by providing a conductive substrate (step 804) with a catalyst metal coating (step 808). FIG. 9A is an enlarged cross-sectional view of part of a conductive substrate 904 with a metal coating 908, forming part of a cathode 912. In this example, the conductive substrate 904 is steel. The metal coating 908 is copper. The conductive substrate may be in the form of a net over which the metal coating is applied. In other embodiments, the conductive substrate and metal coating may be a single piece of the same material, such as a copper wire. In such a case, the metal coating may be considered an outer layer of the metal substrate.
A metal oxide coating is formed on the catalyst metal (step 808). FIG. 9B shows the part of the cathode 912 after the metal oxide coating 916 is formed. In this example, part of the copper catalyst metal coating 908 is formed into copper oxide by heating the cathode to at least 300° C. for at least 15 minutes. Preferably, the metal oxide coating is thicker than a native oxide layer. For example, the metal oxide coating has a thickness of at least twice the thickness of a native metal oxide layer. More preferably, the metal oxide coating is at least 50 nm thick. In other embodiments, the metal oxide coating 916 may be provided by a deposition process to deposit the metal oxide coating on the catalyst metal coating. In this example the copper catalyst metal coating 908 and the metal oxide coating 916 form a metal and metal oxide coating. In other embodiment, metal particles and metal oxide particles may form a single layer to form the metal and metal oxide coating.
In some embodiments, some or all of the native metal oxide layer may be reduced before or during usage as a cathode. In the specification and claims, the term “oxidized cathode” will apply to a cathode on which an oxide layer is formed on the cathode by a process that increases the thickness of the metal oxide beyond that of a native metal oxide, whether the metal oxide coating remains or is subsequently reduced. Therefore the oxidized cathode is a cathode with a oxidized cathode layer, which is a metal and metal oxide coating where the metal oxide either remains or is reduced back to metal, and wherein the metal oxide is at least twice as thick as native metal oxide.
An anode is spaced apart from the cathode (step 812). FIG. 10 is a schematic view of an electrolyzer 1000 that may be used in an embodiment of the invention. An anode is formed by a conductive anode substrate 1004 covered with an anode material 1008. In this example, the anode material 1008 is nickel. An anode electrolyte compartment 1012 is adjacent to the anode and holds an anode electrolyte. The anode electrolyte is provided from an anode electrolyte source 1016, which may continuously circulate anode electrolyte through the anode electrolyte compartment 1012. A cathode electrolyte compartment 1020 holds a cathode electrolyte. The cathode electrolyte is provided from a cathode electrolyte source 1024, which may continuously circulate cathode electrolyte through the cathode electrolyte compartment 1020. Alternatively, the cathode electrolyte may flow to a tank where the solution-phase products are collected. A separator 1028 is placed between the anode electrolyte compartment 1012 and the cathode electrolyte compartment 1020. The separator 1028 may be a porous frit or membrane that may allow certain ions to pass through the separator 1028. As described above, a cathode comprising a conductive substrate 904 with an oxidized cathode layer 1032 forms a cathode adjacent to the cathode electrolyte compartment 1020. In this embodiment, a gas chamber 1036 is placed on the backside of the cathode. A CO2 source 1040 provides a flow of CO2 into the gas chamber 1036. A product collector 1044 collects gas-phase products and unused CO2 from the gas chamber 1036. Product in the product collector 1044 may be isolated and the remaining CO2 may be recycled back to the CO2 source 1040. A voltage source 1048, such as a battery, provides a voltage between the anode and cathode.
In operation, the anode electrolyte source 1016 flows electrolyte through the anode electrolyte compartment 1012. The cathode electrolyte source 1024 flows electrolyte through the cathode electrolyte compartment 1020. CO2 is flowed from the CO2 source 1040 into the gas chamber 1036. The voltage source 1048 applies a positive voltage to the anode substrate 1004 and a negative voltage to the cathode substrate 904 with the anode connected to a positive terminal and the cathode connected to a negative terminal. The process provides electrolysis of the CO2. Various chemical reactions may occur during the electrolysis of CO2, depending on the metal cathode and other factors. One chemical reaction is CO2+H2O→CO+H2+O2. Other chemical reactions provide products of HCO2H, CH3OH or C2H4. In a preferred embodiment, the product collector 1044 provides the product to another system that converts CO, O2, and H2 and possibly other products to methanol or some other fuel or usable chemical.
It has been unexpectedly found that by providing a metal oxide layer on a cathode that is thicker than the native oxide layer and subsequently reducing the metal oxide layer, the reduction of CO2 is improved. Without being bound by theory, it is believed that the reduction of the thick metal oxide layer results in metal particles that have unique structures that result in improved CO2 reduction, however, the reason for the improvement is currently unknown. It has also been unexpectedly found that for some cathodes having a metal and metal oxide interface improves CO2 reduction. Preferably, the metal and metal oxide use the same metal material. However, in an embodiment using cerium oxide, the metal is something other than cerium such as tin or copper. Since cerium would turn to cerium oxide during electrolysis, tin is used to provide a native metal for an enhanced metal oxide metal interface, which provides improved CO2 reduction.
As demonstrated above, copper cathodes that are annealed at 130° C. to grow the oxidation layer do not provide the desired improvement. Annealing copper at 300° C. provides some improvement. It has been found that annealing copper at over 500° C. provides the preferred improvement. Anodization at a constant potential for several hours can also be used to obtain a thick Cu2O layer on Cu and results in improved CO2 reduction.
In the case of some metals such as gold, neither annealing in air or O2 or anodization at a constant potential is effective for preparing a thick oxide layer. Instead, a square wave potential routine is preferred to obtain the metal oxide layer. In the case of gold, a thick, hydrous Au2O3 layer can be formed on the Au electrode by applying a square wave potential alternating between 2.7 V and 0.45 V vs Hg/HgSO4 at a frequency of 1 kHz for 30-60 min. Subsequent reduction of this Au2O3 layer results in a Au electrode with greatly improved CO2 reduction activity and resistance to catalyst deactivation. Similarly, growth of a silver oxide on silver electrodes by application of a square wave potential routine, followed by electrochemical reduction, results in a Ag electrode with greatly improved CO2 reduction activity and resistance to catalyst deactivation.
While this invention has been described in terms of several preferred embodiments, there are alterations, permutations, modifications and various substitute equivalents, which fall within the scope of this invention. It should also be noted that there are many alternative ways of implementing the methods and apparatuses of the present invention. It is therefore intended that the following appended claims be interpreted as including all such alterations, permutations, modifications, and various substitute equivalents as fall within the true spirit and scope of the present invention.

## Claims

1. A method for electrochemically reducing CO2 comprising:
providing a cathode, wherein the cathode comprises a conductive substrate with a catalyst of a metal and a metal oxide based coating on a side of the cathode, wherein the providing the cathode comprises:
providing a conductive substrate with a metal coating;
providing on the conductive substrate a metal oxide coating by either annealing the metal coating, electrochemically oxidizing the metal coating, chemically oxidizing the metal coating, or depositing a metal oxide layer; and
reducing metal oxide in the metal and metal oxide based coating to the metal 0 oxidation state;

providing an anode spaced apart from the cathode;
providing an ionic transport between the anode and cathode;
exposing the cathode to CO2 and H2O
exposing the anode to H2O;
providing a voltage between the cathode and anode; and
reducing CO2 to CO and/or HCO2H.

2. The method, as recited in claim 1, wherein the metal oxide is tin oxide, copper oxide, silver oxide, palladium oxide, gold oxide, molybdenum oxide, lead oxide, platinum oxide, nickel oxide, bismuth oxide, antimony oxide or cerium oxide.

3. The method, as recited in claim 2, wherein the metal oxide is thicker than a native oxide.

4. The method, as recited in claim 3, where the metal oxide in the metal and metal oxide based coating has a thickness that is greater than 50 nm.

5. The method, as recited in claim 1, where the metal oxide in the metal and metal oxide based coating has a thickness that is greater than twice a thickness of a native oxide layer.

6. The method, as recited in claim 1, wherein the metal oxide and metal are of the same metal material.

7. The method, as recited in claim 1, wherein the metal and metal oxide containing coating provide a metal and metal oxide interface.

8. The method, as recited in claim 1, wherein the exposing the cathode to CO2 and H2O, comprises:
exposing a first side of the cathode to H2O; and
flowing CO2 past a second side of the cathode.

9. The method, as recited in claim 1, wherein the providing the cathode, comprises:
applying an anodic square wave potential to the metal coating to form an oxide layer.

10. The method, as recited in claim 9, wherein the metal coating is gold or silver.

11. The method, as recited in claim 1, wherein the cathode is copper and the coating is formed by heating the cathode to a temperature of at least 500° C. for at least 15 minutes.

12. The method, as recited in claim 1, wherein the cathode is copper and the coating is formed by heating the cathode to a temperature of at least 300° C. for at least 15 minutes.

13. The method, as recited in claim 1, wherein the reducing CO2 to CO and/or HCO2H reduces CO2 to HCO2H.