Probing the Mechanism of Aqueous CO2 Reduction on Post

Subscriber access provided by University of Akron

Article 2

Probing the Mechanism of Aqueous CO Reduction on PostTransition Metal Electrodes using ATR-IR Spectroelectrochemistry James E. Pander, Maor F. Baruch, and Andrew B. Bocarsly ACS Catal., Just Accepted Manuscript • DOI: 10.1021/acscatal.6b01879 • Publication Date (Web): 10 Oct 2016 Downloaded from http://pubs.acs.org on October 12, 2016

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a free service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are accessible to all readers and citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

ACS Catalysis is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis

Probing the Mechanism of Aqueous CO2 Reduction on Post-Transition Metal Electrodes using ATR-IR Spectroelectrochemistry James E. Pander III, Maor F. Baruch, and Andrew B. Bocarsly* Department of Chemistry, Princeton University, Princeton, New Jersey 08544, United States Keywords: CO2 reduction, CO2 utilization, in situ IR spectroscopy, p-block metals, spectroelectrochemistry, electrocatalysis

Abstract: The role of metastable surface oxides in the reduction of CO2 on lead, bismuth, tin, and indium electrodes was probed using in situ attenuated total reflectance infrared (ATR-IR) spectroelectrochemistry. The impact of the surface oxide on the Faradaic efficiency of CO2 reduction to formic acid was studied by etching and anodizing the electrodes, and the results were correlated with respect to the observed spectroscopic behavior of the catalysts. A metastable oxide is observed on lead, tin, and indium cathodes under the electrochemical conditions necessary for CO2 reduction. Spectroscopic evidence suggests that bismuth electrodes are fully reduced to the metal under the same conditions. The dynamics of the electroreduction of CO2 at lead and bismuth electrodes appears to be different than on tin and indium electrodes, which suggests that these catalysts act through different mechanistic pathways. The posttransition metal block can be divided into three classes of materials: oxide active materials, oxide

1 ACS Paragon Plus Environment

ACS Catalysis

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 2 of 27

buffered materials and oxide independent materials, and the mechanistic differences are discussed.

1. INTRODUCTION In the past century, a drastic increase in the consumption of fossil fuels has led to elevated levels of carbon dioxide in the atmosphere.1 Such an increase has been associated with negative environmental effects such as an increase in sea level, perturbations to the water cycle impacting agriculture, and an increase in transmission of some communicable diseases due to higher insect populations.2 The necessity of mitigating such effects has recently sparked an increased interest in reducing the atmospheric partial pressure of carbon dioxide via capture and utilization.3–5 The electrochemical reduction of CO2 to value-added products is a promising method for moderating the atmospheric concentrations of CO2. As such, there has been considerable recent interest in the development cathode materials for the reduction of CO2. Many materials have been studied to determine their viability for this application.4 The overwhelming majority of metals are not catalytic for CO2 reduction, with a few exceptions. These exceptions can be classified by their primary CO2 reduction product. Materials like gold and silver primarily produce CO, the post transition metal block primarily produces formate, and copper produces hydrocarbon products.4,5 However, the commercial use of such systems is hindered by their low selectivity over the hydrogen evolution reaction (HER), low rates of product formation, high overpotentials, and relatively fast cathode deactivation.4,6 Such shortcomings can be partially mitigated by a proper understanding of the mechanistic factors that govern the reduction of CO2 on metal electrodes. Presently, mechanistic insight into the reduction pathways is limited, however.


Page 3 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis

Recent work has demonstrated that the catalytic capabilities of many of these materials exhibit a strong dependence on the amount of metal oxide present at the electrode/electrolyte interface.7– 11

This effect is particularly well-documented for the p-block metals In, Sn, and Pb. It has been

previously shown that removing the surface oxide on tin electrodes decreases the Faradaic yield of reduced products.7 Likewise, an increase in the amount of surface oxide on either indium or tin electrodes causes a significant improvement in their Faradaic yields of formic acid and carbon monoxide.7,9 A recent report by Lee and Kanan claimed that oxide-derived lead electrodes demonstrate high Faradaic yields of CO2 reduction products by suppressing the hydrogen evolution reaction.8 Although tin,7,11–14 indium,9,14–16 and lead4,8,14,17,18 electrodes have been widely studied as CO2 reduction catalysts, there have been relatively few reports on the use of bismuth electrodes for CO2 reduction in aqueous media.19,20 To date, both morphological and chemical effects have been put forth as root causes for the apparent oxide-dependence of CO2 reduction catalysts,7–9,21–24 although such effects are not fully understood. Our research group recently reported IR spectroscopic studies of the electrode/electrolyte interface of tin and indium electrodes, demonstrating the presence of a surface-bound carbonate species formed via the reaction of CO2 with a surface hydroxide, that is a key intermediate in the reduction of CO2 at these interfaces.9,10 Here, we consider the role of surface oxides in the reduction of CO2 at lead and bismuth electrodes. 2.

EXPERIMENTAL METHODS

2.1.

Electrochemical Measurements

All electrochemical experiments were performed with a standard three electrode setup with an aqueous Ag/AgCl (3 M NaCl) reference electrode and a platinum mesh (Sigma Aldrich, 99.9%) auxiliary electrode. All potentials are reported against this reference. For electrolysis


ACS Catalysis

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 4 of 27

experiments, the auxiliary electrode was encased in a gas dispersion tube (Ace Glass, 4-8 μm porosity) to separate the two electrodes and prevent the reoxidation of CO2 reduction products. The working electrodes for cyclic voltammetric and electrolysis experiments consisted of 1 cm x 1 cm flags of bismuth (Goodfellow, 1.00 mm thick, 99.9999%), lead (Alfa Aesar 0.25 mm thick, 99.998%), tin (Sigma Aldrich, 0.5 mm thick, 99.998%), and indium (Sigma Aldrich, 0.25 mm thick, 99.999%) connected to copper wire with conductive silver epoxy (Epoxy Technology E3037) and encased in a glass tube with insulating epoxy (Loctite 0151). Experiments were conducted in a continuously stirred aqueous 0.1 M K2SO4 (Sigma Aldrich, > 99%) electrolyte and purged with either argon or carbon dioxide (Airgas). Electrolysis experiments were conducted until 4 C of charge had passed (typically 15-20 minutes). Electrochemical measurements were collected with CH Instruments 760D and 1140B potentiostat/galvanostat workstations. Liquid electrolysis products were detected using a Bruker 500 MHz 1H NMR with a cryoprobe detector and a custom solvent suppression pulse sequence. Formic acid calibration curves were generated using a stock solution of formic acid (J. T. Baker, 88%) using 1,4-dioxane as an internal standard. Once the amount of products was detected from these calibration curves, faradaic yields were calculated by taking the ratio of charge passed to form the product of interest to the total charge passed. In the cases of anodized electrodes, a large current corresponding to the initial reduction of the bulk oxide is observed before the current decays to a steady state. This generally occurs very quickly; however, in the case of lead, it can take up to several minutes and account for a non-negligable amount of the total current passed during the experiment. It was determined via 1H NMR that no formate is produced during this initial phase


Page 5 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis

of the electrolysis, so the charge passed during this portion of the electrolysis was not included in the faradic yield calculations. In order to control the surface oxidation of the catalysts, the surfaces were either acid-etched to remove the oxide or electrochemically anodized to provide a thick oxide layer. Electrodes were etched through immersion in a solution of boiling HBr (24%, Sigma Aldrich) for several minutes, rinsed with DI water, and transferred to the electrochemical cell immediately to minimize reoxidation. The anodization procedure for each type of electrode varied slightly. Bismuth was held at +3 V vs Ag/AgCl in 0.025 M H2SO4 for 3 minutes,25 lead was held at +3 V vs Ag/AgCl in 1 M H2SO4 for 3 min,26 tin was held at +5 V vs Ag/AgCl in 0.5 M oxalic acid for 3 min,27 and indium was held at +3 V vs Ag/AgCl in 1 M KOH for 3 min.9 2.2.

Thin Film Preparation and ATR-IR Spectroscopy

ZnSe ATR-IR crystals were modified with a 400 nm thick layer of metal using ion beam sputtering as was previously described.10 The conductivity of each sample was tested across a ~1 cm portion of the layer, with the In, Sn, and Pb layers having < 0.1 Ω resistance and the Bi layers having < 50.0 Ω resistance. Additionally, due to the fragility of pure Bi thin films, a 25 nm carbon underlayer was deposited underneath the 375 nm of Bi. The carbon underlayer helped prevent the Bi layer from delaminating under an applied potential, though the film was still susceptible to mechanical stress. The films were then utilized for direct in situ electrochemical ATR-IR experiments as described in our previous work.10 2.3. Surface Characterization X-ray photoelectron spectroscopy (XPS) experiments were conducted on a VG Scientific MK II ESCALab with a magnesium X-ray source. The detector was a hemispherical analyzer with a 20 eV pass energy. Spectra were deconvoluted and quantitated using the CasaXPS software


ACS Catalysis

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 6 of 27

package. Scanning electron microscopy (SEM) images were collected using a Quanta 200 FEG ESEM instrument. 3.

RESULTS AND DISCUSSION

3.1.

Electrochemical Behavior of Lead and Bismuth Electrodes

The cyclic voltammetric responses of lead and bismuth electrodes at a scan rate of 50 mV/s in pH 4.4 Ar- and CO2-saturated 0.1 M K2SO4 are shown in Figure 1. The lead electrode exhibits a set of waves centered at -0.54 V vs Ag/AgCl with a separation of 270 mV. The peak to peak separation is invariant with change in scan rate. Peaks at this position have been previously correlated with the reduction of PbO to Pb0 and its subsequent reoxidation.28–30 Other than the anticipated increase in current with scan rate, there is no notable variance in the cyclic voltammograms between 10 mV/s and 1 V/s. Upon exposure to CO2, there is little change in this redox feature; however, at significantly more negative potentials an increase in current is observed starting around -1.5 V vs Ag/AgCl which has previously been associated with the reduction of CO2 to formate.4 Similarly, bismuth electrodes exhibit a redox couple with a halfwave potential of -0.50 V vs Ag/AgCl and a separation of 270 mV at a scan rate of 50 mV/s that has been associated with the Bi2O3/Bi0 transformation.31,32 A small anodic wave at approximately -0.075 V corresponding to the oxidation of Bi0 to Bi3+ is also observed. In contrast to lead, when exposed to CO2, this process becomes irreversible (highlighted in the Figure 1b inset) and the reduction portion of the wave shifts to slightly more positive potentials. Additionally, current enhancement starting around -1.5 V vs Ag/AgCl is observed corresponding to CO2 reduction to formic acid.4,19 Similar surface oxide redox features have been noted on indium and tin electrodes, with each dramatically diminishing in magnitude when exposed to CO2 in contrast to the behaviors of Pb and Bi noted here.9,10


Page 7 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis

In order to further evaluate the reduction event observed on all of these materials, Figure 2 shows the Faradaic efficiency for the production of formate on indium, tin, lead, and bismuth electrodes as a function of their surface oxidation in the low overpotential regime for each material. Indium and tin electrodes exhibit a strong dependence on the oxidation of the electrode surface, as has been previously reported.7,9,10 The efficiency for formate production decreases when the electrodes are chemically etched, removing the native oxide, and increases when the electrode is modified with a thick oxide layer via anodization.7,9 However, the effects of oxidation on lead and bismuth electrodes appear to be more nuanced. The Faradaic yield of formate on lead electrodes increases significantly when the amount of oxide is increased via anodization, compared to the native Pb surface. Likewise, when the native lead surface oxide is removed by chemical etching an increase in Faradaic yield is observed. Bismuth electrodes with native surfaces reduce CO2 to formate with high efficiency (~50%) compared to the other native surfaces of the metals under consideration. There is no statistically significant change in Faradaic efficiency upon anodization, indicating that there is no significant benefit from increasing the amount of surface oxide. When etched, the Faradaic efficiency decreases, which initially suggests that the oxide may play a role. However, if exposed to air to allow the oxide to regrow (as determined by XPS), the Faradaic yield for formate production does not return to that of the native foil as it does in the case of indium and tin electrodes, indicating that the reason for the decrease in efficiency is more subtle. Figure 3 shows a cyclic voltammetric comparison between native and etched electrodes of both lead and bismuth. When a bismuth electrode is etched, there is a significant decrease in the overpotential for the HER, shown by the shift in current onset to lower potentials. In contrast, on lead electrodes, while there is slight increase in current, the current onset for proton reduction does not change. The etching


ACS Catalysis

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 8 of 27

procedure affects the two materials differently; it is possible that, in addition to chemically removing the oxide, the etchant may alter the physical or chemical structure of the electrode. In order further probe this possibility, scanning electron microscopy (SEM) images of each electrode material were obtained and compared as described in Section 3.2. The partial current densities of formate and hydrogen were collected as a function of pH on native lead and bismuth electrodes at -1.9 V vs Ag/AgCl in Ar and CO2 saturated 0.1 M K2SO4, as shown in Figure 4. Current densities are calculated from potentiostatic bulk electrolysis experiments as the product of the calculated Faradaic efficiency for each product in that experiment and the average current density of the experiment. At these potentials, the rate of CO2 reduction should be at a maximum based on the peak currents of cyclic voltammetric experiments. The partial current density for H2 in the Ar saturated solutions gives a baseline of the capability of each electrode to produce H2. In the case of lead cathodes, when CO2 is introduced to the system, there is little to no change in the electrode’s ability to reduce protons to H2, as indicated by the invariant partial current density for H2 between the Ar and CO2 saturated cases. The electrode continues to produce H2 at the same rate, while producing formate in addition. This suggests that the reactions to form H2 and formate are not directly competing for surface sites. In contrast, on bismuth cathodes, when CO2 is introduced, the partial current density for H2 formation is significantly suppressed when compared to the Ar saturated case. Thus the formation of formate is directly competing with the formation of H2, suggesting that these two processes are competing for surface reaction sites. 3.2.

Surface characterization (SEM and XPS) of Pb and Bi Foils

In order to probe the effects that varying the amount of surface oxide plays on electrode morphology, SEM images of native, etched, and anodized samples (before and after electrolysis)


Page 9 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis

of each of the four p-block materials were obtained (Figure 5). In each case, the native sample is relatively smooth, and treating these samples causes different changes to the surface morphologies for each material. Chemical etching causes significant morphological changes to bismuth and tin surfaces. On bismuth, etching causes roughening of the surface through the formation of a large number of convex microstructures. However, the tin surface becomes significantly pitted, with pore sizes on the micron scale. The surface roughness of the lead and indium samples does not appreciably change upon etching. Bromine was observed on the surface of the etched samples via XPS; however glancing incidence X-ray powder diffraction did not reveal metal bromide phases at the sample surface, indicating that the amount of bromine present after etching is small and localized to the immediate surface region. Anodization has a different effect on surface morphology. In particular, the lead oxide and bismuth oxide surfaces do not appear to be appreciably rougher than those of the native oxide. However, this is not the case for the tin and indium samples, on which increased microstructuring is observed on the anodized samples. Notably, anodized electrodes undergo significant morphological changes during the course of an electrolysis experiment, most prominently observed in the case of the post-anodized lead electrode, which goes from being relatively smooth to exhibiting a very rough surface morphology. X-ray photoelectron spectroscopy experiments were conducted to gain insight into the oxide dependence on lead and bismuth. The Pb 4f region is shown on the left of Figure 6, while the Bi 4f region is shown on the right. Each chemical species is presented by a pair of peaks due to the effects of spin orbit coupling, which correspond to the 4f7/2 and 4f5/2 spin states. The signal from the native lead foil is well fit as a single species with binding energies of 138.9 eV and 143.8 eV for the 4f7/2 and 4f5/2 spin states, respectively. This agrees well with previous reports of a native


ACS Catalysis

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 10 of 27

PbO species (trace a). The lack of peaks corresponding to lead in the metallic state indicates that the native lead oxide is thicker than the electron escape depth, several nanometers. Upon anodization, this peak broadens and shifts to lower binding energy, as shown in trace b. Deconvolution reveals two chemical species with the first occurring at binding energies of 138.4 eV and 143.3 eV corresponding to Pb2+and with the second species occurring at binding energies of 137.6 eV and 142.5 eV corresponding to Pb4+. These peaks match well with previously reported Pb2+ and Pb4+ signals from a mixed valent Pb3O4 structure.33 Thus, not only does anodization thicken the oxide, it also chemically alters the surface. Alternatively, when the lead sample is etched in HBr, two sets of well resolved peaks are observed with the set at 138.8 eV and 143.7 eV again corresponding to PbO and the set at 136.6 eV and 141.5 eV corresponding to Pb0.33 The oxidation of the surface is likely due to the quick reoxidation of the metallic surface upon exposure to atmospheric oxygen during transfer from the etching solution to the spectrometer. When an anodized electrode is electrochemically reduced at -1.6 V vs Ag/AgCl (trace d), the metallic lead peak grows substantially, indicating that the thick oxide film grown during the anodization has been significantly reduced. PbO is also observed at the reduced electrode surface. The bismuth samples behave in a similar way. Again, the signal corresponding to the 4f7/2 and 4f5/2 states containing identical chemical information is observed for each sample. The signal from a native bismuth foil (trace a) can be deconvoluted to two chemical species, with the major chemical species at binding energies of 159.2 eV and 164.5 eV and the minor chemical species at a binding energy of 156.9 eV and 162.2 eV corresponding to Bi2O3 and Bi0, respectively.34 Thus, the surface of bismuth is natively oxidized, though to a lesser degree than lead. When anodized (trace b), the peaks corresponding to the metallic bismuth disappear, leaving only the bismuth


Page 11 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis

oxide peaks at 159.6 eV and 164.9 eV. This is consistent with a thickening of the oxide film such that the oxide is now thicker than the escape depth of the electrons. When bismuth is etched in HBr, the oxide peaks diminish almost into the baseline, while the metallic peaks at 156.8 eV and 162.1 eV become the dominant feature, as shown in trace c. The fact that the Bi2O3 peaks are small indicates that very little of the surface reoxidizes upon transfer of the sample to the spectrometer, which may suggest that the etched bismuth surface oxidizes more slowly than that of the etched lead surface. Electrochemical reduction of the anodized bismuth electrode at -1.5 V vs Ag/AgCl gives XPS peaks originating from both Bi2O3 and Bi0 and indicating the reduction of the bismuth oxide under electrochemical conditions. 3.3. In Situ ATR-IR Spectroelectrochemistry of Pb and Bi Interfaces Indium and tin electrodes have previously been demonstrated to exhibit a kinetically metastable oxide at the electrode/electrolyte interface at electrode potentials where CO2 reduction is observed.8,9,10 At these potentials, the thermodynamically favored state is a purely metallic surface.35 It has generally been assumed that reduction of CO2 on metal cathodes involves the initial formation of CO2•- at the bare metal surface, and both inner and outer sphere pathways have been invoked.4 However, our previous work on indium and tin electrodes has shown that it is the metastable oxide on these electrode interfaces instead of the bare metallic surface that interacts with CO2.9,10 Thus, careful characterization of the surface under applied potentials is necessary to understanding the true state of the chemically active interface of these systems. In situ ATR-IR spectroelectrochemistry has proven to be useful in this endeavor. In our previous work, it has been shown that it is possible to obtain information about the relative thickness of metastable surface oxides as well as the chemical nature of the oxide. Under applied potential, vibrational stretches at 3500 cm-1 and 1650 cm-1 (corresponding to water) emerge in the


ACS Catalysis

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 12 of 27

ATR-IR spectra due to the oxide layer physically thinning as it is reduced.9,10 With a thinner oxide layer, more water enters the static sampling region of the evanescent IR beam, causing enhanced absorbance at the indicated vibrational frequencies. This behavior was shown to be reversible on indium and tin cathodes. Specifically, the oxide layer could be thinned or thickened depending on the applied potential, causing the intensity of the water vibrations to vary inversely with the oxide thickness.9,10 To probe this behavior on lead and bismuth, similar experiments were conducted. In these experiments, a background spectrum was taken under open circuit conditions before applying a potential of -1.2 V vs. Ag/AgCl, which is more negative than the metal oxide-to-metal reduction wave observed in cyclic voltammetry. Upon applying the potential, a large increase in absorbance at 3500 cm-1 and 1650 cm-1 was observed (Figure 7). In order to test the reversibility of this change, a new background was taken and then a potential more positive than the cyclic voltammetric redox wave was applied (0.0 V vs. Ag/AgCl). This potential jump generated an anodic current. An ATR-IR spectrum was then taken, as shown in trace b, of Figure 7. Again a new background scan was taken, before the potential was reversed to -1.2 V, reducing the surface oxide again. The spectra observed after this reversal are shown in trace c. Both bismuth and lead exhibit an initial reduction of the native oxide leading to the strong water signal observed in trace a. On lead cathodes, when the potential was jumped from -1.2 V vs Ag/AgCl to 0.0 V vs Ag/AgCl, the signals at 3500 cm-1 and 1650 cm-1 became negative, indicating the exclusion of water due to a thickening of the oxide at the electrochemical interface. When the electrode was returned to -1.2 V vs Ag/AgCl, these peaks return to having positive intensity. In the case of bismuth, after the initial reduction, jumping the potential positive, induced a very slow reoxidation of the electrode surface, as evident by the very small


Page 13 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis

current passed and the very low intensity of the associated water signal in the IR. Returning the electrode to -1.2 V vs Ag/AgCl, caused these peaks to become positive again. These experiments demonstrate that lead oxides are slowly but reversibly reduced at the electrode surface, while bismuth oxides reduce quickly but oxidize slowly, and are thus semi-reversible. In order to probe this further, the intensity of the included water signal was observed as a function of applied potential in 50 mV steps. Figure 8a and 8b show the spectroscopic behavior of lead and bismuth electrodes, respectively, as the potential is gradually decreased in an Arsaturated electrolyte. A background was taken under open circuit conditions before a potential of -1.0 V vs. Ag/AgCl was applied to the electrode and the change in the ATR-IR difference spectrum was recorded. As in the previous experiment there is initially a large inclusion of water immediately upon the application of a potential. In order to observe small changes in the oxide layer as the potential was changed, a new background spectrum was taken at this potential before stepping the potential more negative by 50 mV. The electrode was allowed to equilibrate at each potential for 5 minutes before being stepped to the next potential. This process was repeated until the low CO2 overpotential regime for each electrode was reached (-1.6 V vs Ag/AgCl for lead and -1.5 V vs Ag/AgCl for bismuth). The differences in behavior of the oxide at the electrode surface of lead and bismuth electrodes under operating conditions are striking. On lead cathodes, after the initial reduction of the native oxide and as the potential is stepped cathodically, the IR signals due to the inclusion of water are continually observed, indicating that the oxide is continuing to thin. This demonstrates that a metastable lead oxide is present at potentials negative enough to reduce CO2 on lead electrodes. In contrast, on bismuth electrodes, after the initial oxide layer thinning, there is very little change in the ATR-IR water related frequencies as the potential is changed, indicating that there is little


ACS Catalysis

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 14 of 27

change to the electrode surface over the experiment. We therefore conclude that the native bismuth oxide reduces to an invariant state at modest potentials. Based on the spectroscopic data it is not possible to determine if the oxide is no longer present, or present but unchanging. In order to probe the interaction of the CO2 with the electrode surface, CO2 was introduced into the electrochemical cell. The results are shown in Figure 9. In order to determine what is occurring at the electrode surface during the CO2 reduction reaction, a background IR spectrum was taken under an argon atmosphere with the electrode at the operating potential, then data was collected after bubbling CO2 into the electrochemical cell. When CO2 is introduced into the electrochemical cell, peaks centered at 2350 cm-1 corresponding to gaseous CO2 are observed. As has been previously reported, on indium and tin cathodes, peaks at 1500 cm-1 and 1385 cm-1 corresponding to surface confined metal-carbonate species are observed.9,10 These peaks shift with 13C isotopic labelling, as is expected. In contrast, on lead and bismuth cathodes, no additional peaks are observed in this region. Lead carbonate is reported to yield a single absorption at 1410 cm-1.36 Similarly, bismuth carbonate is reported to exhibit a single very strong absorption at 1389 cm-1.37 The lack of signal corresponding to an interfacial carbonate on lead and bismuth has significant implications on the mechanism of CO2 reduction on these materials. Based on this observation, it is likely that CO2 is directly reduced at the metallic sites of lead and bismuth electrodes. 3.4. Mechanistic Discussion Careful consideration of the data allows us to propose mechanisms for the CO2 reduction on lead and bismuth cathodes. As shown by the in situ ATR-IR data shown in Figure 9, the lack of


Page 15 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis

the observed surface carbonate intermediate suggests that lead and bismuth reduce CO2 to formate through entirely different mechanisms compared to indium and tin. Using the same techniques that have successfully identified metastable oxide species on tin and indium, the spectroscopic data presented in Figures 7 and 8 clearly demonstrate that a similar metastable oxide is present on the surface of lead cathodes during the reduction of CO2. This observation has also been previously reported by Lee et al.8 Despite the presence of a metastable oxide under relevant conditions, CO2 does not interact with this oxide to form a spectroscopically observable carbonate as it does on tin and indium. However, additional information is needed to determine if the reaction occurs at a PbO site or a Pb0 site. When lead electrodes are etched in boiling HBr to remove the surface oxide, the faradaic yield for formate production significantly increases. Simultaneously, as demonstrated by XPS, the amount of Pb0 sites at the surface increase. Notably, the etching procedure does not completely remove the PbO on the electrode surface, as there is still PbO observed by XPS. On the other hand, anodized electrodes demonstrate an increased ability to reduce CO2 to formate as well. Lee et al. have claimed that the increase in yield upon anodization is due to a suppression of the competing hydrogen evolution reaction instead of an increase in reactivity towards CO2.8 The XPS data presented in Figure 5 shows that upon anodization, lead electrodes grow a thick, mixed valent Pb3O4 species and that after being used as the cathode in a CO2 electrolysis experiment, a mix of PbO and Pb0 are observed by XPS. This indicates that during the course of the electrolysis experiment, the lead oxide species is reduced from a thick Pb3O4 to a thin or porous PbO species. Notably, this change occurs within the first few minutes in the electrolysis experiment, meaning that throughout the course of the CO2 reduction reaction, the surface is a mix of PbO and Pb0 sites. Additionally, when the anodized electrode undergoes this change in chemical speciation,


ACS Catalysis

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 16 of 27

the morphology of the surface is significantly changing, more so than when any other anodized p-block material is used in an electrolysis experiment, as shown in Figure 5. These chemical and morphological changes clearly yield a more catalytically active surface, and combined with the spectroscopic data, suggest that the surface oxide is an important factor in the CO2 reduction mechanism, but that it does not play a direct chemical role in transferring charge to CO2. Interestingly, there is little change in the partial current density for formate production between pH 4.4 and 6.5 on native lead surfaces. This suggests that the rate determining step is independent of the bulk concentration of protons, which may be due to an internal buffering effect that occurs at the electrode surface. Additionally, we note that the partial current density for H2 formation remains unchanged between Ar saturated solutions and CO2 saturated solutions. This indicates that the CO2 reduction reaction is not competing with the H+ reduction reaction for reaction sites at the electrode surface. This finding again supports the idea that the interfacial pH is invariant with the solution pH. Thus it is proposed that the first mechanistic step involves an acid/base reaction between solution protons and the electrode surface. The pKa of the surface oxide is reported to be approximately 7.7.38 Notably, when electrolysis experiments are performed at pH values higher than this pH, no formate production is observed, supporting the notion that protonation of the surface is important. Many metals, including lead electrodes are known to reduce protons to adsorbed hydrogen.39 The self-exchange current has been reported to be around -10-9 A/cm2 and the chemisorption bond strength has been determined to be approximately 125 kJ/mol.40 Thus, once the oxide at the surface of lead has been protonated, it is likely that the first charge transfer step involves the reduction of this proton to an adsorbed hydrogen atom. This adsorbed hydrogen atom can then react with other species.


Page 17 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis

The next reaction step likely involves the coupling of the surface hydrogen with either another surface hydrogen to form H2 gas through the Volmer reaction or with CO2 and a proton to form formate. Based on the observation that Pb0 sites play an important role in the reduction of CO2 and the apparent lack of interaction between CO2 and the oxide in the IR experiments, it is likely that CO2 reduction occurs at a metallic site, while the formation of hydrogen occurs at the oxide. Thus the overall mechanism can be summarized as: (1) (2) (3) (4) where the rate determining step is either reaction (2) or (4). The data presented do not indicate if the reduction of CO2 occurs as a concerted step with the coupling of the surface hydrogen or as independent steps. On bismuth cathodes, the spectroscopic results presented demonstrate that the oxide does not play a role in the CO2 reduction reaction. Specifically, there is no change in the amount of oxide on the surface as the potential is lowered after an initial reduction, as shown by Figure 8, and the reaction does not proceed through an observable intermediate, as shown in Figure 9. This is further supported by the observation that the faradaic yields for formate formation are largely independent of the amount of surface oxide. In addition, we observe a large decrease in the partial current density for H2 formation when comparing Ar and CO2 saturated solutions (as shown in Figure 4), indicating that proton reduction to H2 and CO2 reduction to formate are competing for reaction sites at the electrode surface. We therefore conclude that both the


ACS Catalysis

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 18 of 27

reduction of H+ to H2 and the reduction of CO2 to formate occur at metallic bismuth sites at the electrode surface. Thus, we propose a mechanism of the form: (5) (6) (7) where a proton is first reduced to an adsorbed hydrogen atom at a bare metallic site at the electrode surface, followed by an electrochemical reaction of the surface hydrogen with either another surface hydrogen to form H2 via the Volmer reaction (reaction 6), or with CO2 and a proton to form formate (reaction 7). The self-exchange current on bismuth has been reported to be around -10-7 A/cm2 and the bond strength has been determined to be approximately 167 kJ/mol.40 The self-exchange current on Bi cathodes is approximately two orders of magnitude higher than on lead cathodes, which is likely the reason that bismuth cathodes exhibit higher partial current densities for both H2 evolution and CO2 reduction.

4. CONCLUSIONS The data presented demonstrate that lead cathodes support a metastable oxide under CO2 reduction conditions.8 The lack of lead carbonate signal in the ATR-IR spectra presented here suggests that the reduction of CO2 does not proceed through the formation of a surface carbonate on lead cathodes, as it does on tin and indium cathodes. This conclusion is supported by the increase in Faradaic yield upon revealing more of the Pb0 surface via chemical etching, indicating, instead, that CO2 is likely reduced at the bare metal surface sites. Interestingly, the 18 ACS Paragon Plus Environment

Page 19 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis

total faradaic yield for formate production also increases upon anodization of the electrode surface; however, XPS spectra of an anodized lead electrode post electrolysis, shows a significant amount of metallic lead present at the electrode surface, indicating that when the thick oxide is reduced in situ, Pb0 sites are present at the electrode interface in addition to the metastable lead oxide. The roughness of the electrode changes significantly post electrolysis, indicating that it is likely that the increase in yield is due to an increase in surface area. The partial current density for H2 formation is invariant with the presence of CO2, indicating that the hydrogen evolution reaction is not directly competing for surface sites with the carbon dioxide reduction reaction. Thus, we conclude that proton reduction to H2 occurs at the PbO sites and CO2 reduction to formate likely occurs at Pb0 sites. The data support a similar mechanism for CO2 reduction on bismuth, with a notable difference. As demonstrated by the ATR-IR results in the absence of CO2, bismuth is unique out of the formate-producing p-block materials in that it does not support a metastable oxide under reducing conditions. Thus, it appears that CO2 is reduced directly on the bare metallic surface of bismuth, which is again supported by both the lack of observable reaction intermediates in the ATR-IR and by the lack of oxide dependence in the Faradaic yield. The apparent decrease in efficiency of the bismuth electrodes when etched does not appear to be due to an oxide dependence (as is the case on tin and indium) because when the electrode is allowed to reoxidize in air, the faradaic yield does not improve. The partial current density for H2 formation on bismuth cathodes is highly dependent on the presence of CO2, indicating that CO2 and H+ are competing for reaction sites at the electrode surface. Thus, with no oxide present, we conclude that these reaction must occur at the same Bi0 sites.


ACS Catalysis

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 20 of 27

Thus, it appears that there are three types of post-transition metal electrocatalysts: oxide active materials (In and Sn), oxide buffered materials (Pb), and oxide independent materials (Bi). The oxide active materials have been shown to have a strong dependence on the amount of oxidation at the electrode/electrolyte interface. These cathodes support a kinetically metastable oxide under conditions where the surface would be thermodynamically expected to be purely metallic. This oxide is critical in the CO2 reduction mechanism, as it interacts with CO2 to form a metalcarbonate species, which is reduced to formate. Oxide buffered materials support a meta-stable oxide similar to that of the oxide active materials; however, this oxide does not play a direct role in the reduction of CO2 to formate. Instead, this oxide seems to buffer the electrode/electrolyte interface and provide the necessary protons for the reduction of CO2 to formate through this buffering effect. Finally, oxide independent materials differ in that oxide species do not interact with the observed electrochemistry in any way. As presented, it is not clear if the surface is devoid of oxide or if the oxide is present but invariant; however, it is clear that the oxide plays no role in CO2 reduction or H+ reduction, which compete for reaction sites at the electrode surface. The variety in the mechanisms of these materials, while all still producing the same product, may allow for some interesting synergy in the design of novel materials. It may be possible to take advantage of these metastable oxides to create a material where both oxidized and metallic sites are catalytic for the conversion of CO2 to formic acid, potentially creating materials with very high CO2 reduction efficiencies at modest overpotentials.


Page 21 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis

FIGURES

Figure 1: Cyclic voltammograms of (a) lead and (b) bismuth electrodes at 50 mV/s in 0.1 M K2SO4 at pH 4.4 that has been deaerated with argon (blue) or CO2 (red). The inset on the right plot shows a close-up region of the -0.1 V to -1.1 V region of the bismuth scan.

Figure 2: Results of potentiostatic bulk electrolysis experiments on p-block electrodes as a function of surface condition. Bar graphs represent the Faradaic efficiency for formate on (red) lead, (green) bismuth, (grey) tin, and (blue) indium cathodes. Electrolyses were performed under selected low-overpotential conditions for each material, specifically -1.6 V for lead, -1.5 V for bismuth, and -1.4 V for indium and tin.


ACS Catalysis

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 22 of 27

Figure 3. Cyclic voltammograms of a Pb (left) and Bi (right) native foil (blue) and etched foil (red) under an Ar atmosphere at 50 mV s-1 in pH 4.4 0.1 M K2SO4.

Figure 4. Partial current densities for a Pb cathode (left) and Bi cathode (right) at -1.9V vs Ag/AgCl in 0.1 M K2SO4 calculated from the average current and Faradaic yeilds of potentiostatic bulk electrolysis experiments. The black trace (circles) represent the partial current density of H2 formation in an Ar purged electrolyte. The red traces represent the partial current densities of the products in a CO2 purged electrolyte, specifically the partial current density of H2 formation (squares) and the partial current density of formate formation (diamonds).

Figure 5. SEM images of the p-block catalysts used in this work. (a-d) lead (e-h) bismuth (i-l) tin (m-p) indium. The first column represents the native electrode surface, the second column represents the electrode surface after having been etched in boiling HBr, the third column 22 ACS Paragon Plus Environment

Page 23 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis

represents the electrode surface after anodization, and the fourth column represents the anodized electrode’s surface post electrolysis. Throughout, the scale bars represent 25 µm.

Figure 6: XPS spectra of lead (left) and bismuth (right) 4f region of foils under the oxide conditions studied. a) native oxidation b) anodized c) etched in HBr d) post electrolysis sample of an anodized electrode.

Figure 7. Potential jump experiments on a lead (left) or bismuth (right) thin film electrode. The top graphs represent the results of the ATR-IR difference spectra such that (a) a scan is taken at 23 ACS Paragon Plus Environment

ACS Catalysis

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 24 of 27

1.2 V with a background taken under open circuit conditions (b) a scan is taken at 0V with a background at -1.2 V (c) a scan is taken at -1.2 V with a background at 0 V. The bottom graphs represent the current-time trace of the electrochemical experiment during the collection of the spectroscopic data.

Figure 8: ATR-IR traces taken through (a) a 400 nm thin film of Pb and (b) a 375 nm film of bismuth on a 25 nm carbon underlayer as the potential is stepped more negatively. The bottom trace represents a difference spectrum at -1.0 V with a background under open circuit conditions (scaled down by a factor of 10) and each subsequent spectra represents a scan at a 50 mV more negative potential with a background at the previous potential.

Figure 9: ATR-IR spectra of thin films of (a) indium (b) tin (c) lead (d) bismuth when exposed to CO2 under reducing conditions. Peaks at 3500 cm-1 and 1650 cm-1 correspond to the vibrational modes of water observable as surface oxides thin under applied potential. Peaks at 1500 cm-1 and 1385 cm-1 correspond to surface bound carbonate species. The small features at approximately 2900 cm-1 on the bismuth scans are due to impurities in sample preparation that did not appropriately subtract from the background.


Page 25 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis

AUTHOR INFORMATION Corresponding Author *E-mail: [email protected] ACKNOWLEDGMENTS The authors would like to acknowledge the NSF (Grant CHE-1308652) for financial support for this work. We would thank Mr. Matthew Vallon and Mr. James Park for assistance with the collection of XPS and XRD data, respectively. Additionally, we thank Dr. István Pelczer, Mr. Kenith Conover, and Mr. John Schreiber for their assistance with NMR, thin film deposition, and SEM. Finally, we thank Dr. James White for many fruitful discussions. The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript. ABB is a co-founder and president of the Scientific Advisory Board holding less than a 5% interest in Liquid Light, Inc., a company focused on the electrochemical conversion of carbon dioxide to chemicals. ABBREVIATIONS Attenuated total reflectance infrared, ATR-IR; hydrogen evolution reaction, HER; post transition metal block, p-block; x-ray photoelectron spectroscopy, XPS; scanning electron microscopy, SEM


ACS Catalysis

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 26 of 27

REFERENCES (1)

(2)

(3) (4) (5) (6) (7) (8) (9) (10) (11) (12) (13) (14) (15) (16) (17) (18) (19) (20)

Hartmann, D. L.; Klein Tank, A. M. G.; Rusticucci, M.; Alexander, L.; Broennimann, S.; Abdul-Rahman Charabi, Y.; Dentener, F.; Thorne, P.; Wild, M.; Zhai, P. In Climate Change 2013: The Physical Science Basis. Contribution of Working Group I to the IPCC Fifth Assessment Report; Stocker, T. F., Qin, D., Plattner, G.-K., Tignor, M., Allen, S. K., Boschung, J., Nauels, A., Xia, Y., Bex, V., Midgley, P. M., Eds.; Cambridge University Press: New York, 2014. Kunreither, H.; Gupta, S.; Bosetti, V.; Cooke, R.; Dutt, V.; Ha-Duong, M.; Held, H.; Llanes-Regueiro, J.; Patt, A.; Shittu, E.; Weber, E. In Climate Change 2014: Mitigation of Climate Change. Contribution of Working Group III to the IPCC Fifth Assessment Report; Edenhofer, O., Pichs-Madruga, R., Sokona, Y., Farahani, E., Kadner, S., Seyboth, K., Adler, A., Baum, I., Brunner, S., Eickemeier, P., Kriemann, B., Savolainen, J., Schlömer, S., von Stechow, C., Zwickel, T., Minx, J. C., Eds.; Cambridge University Press: New York, 2014. Stauffer, P. H.; Keating, G. N.; Middleton, R. S.; Viswanathan, H. S.; Berchtold, K. A.; Singh, R. P.; Pawar, R. J.; Mancino, A. Environ. Sci. Technol. 2011, 45, 8597–8604. Hori, Y. In Modern Aspects of Electrochemistry; Vayenas, C. G., White, R. E., GamboaAldeco, M. E., Eds.; Modern Aspects of Electrochemistry; Springer New York, 2008; pp 89–189. White, J. L.; Baruch, M. F.; Pander III, J. E.; Hu, Y.; Fortmeyer, I. C.; Park, J. E.; Zhang, T.; Liao, K.; Gu, J.; Yan, Y.; Shaw, T. W.; Abelev, E.; Bocarsly, A. B. Chem. Rev. 2015. Hori, Y.; Konishi, H.; Futamura, T.; Murata, A.; Koga, O.; Sakurai, H.; Oguma, K. Electrochimica Acta 2005, 50, 5354–5369. Chen, Y.; Kanan, M. W. J. Am. Chem. Soc. 2012, 134, 1986–1989. Lee, C. H.; Kanan, M. W. ACS Catal. 2015, 5, 465–469. Detweiler, Z. M.; White, J. L.; Bernasek, S. L.; Bocarsly, A. B. Langmuir 2014, 30, 7593– 7600. Baruch, M. F.; Pander, J. E.; White, J. L.; Bocarsly, A. B. ACS Catal. 2015, 5, 3148–3156. Won, D. H.; Choi, C. H.; Chung, J.; Chung, M. W.; Kim, E.-H.; Woo, S. I. ChemSusChem 2015, 8, 3092–3098. Anawati; Frankel, G. S.; Agarwal, A.; Sridhar, N. Electrochimica Acta 2014, 133, 188– 196. Bumroongsakulsawat, P.; Kelsall, G. H. Electrochimica Acta 2014, 141, 216–225. Mahmood, M. N.; Masheder, D.; Harty, C. J. J. Appl. Electrochem. 1987, 17, 1159–1170. White, J. L.; Herb, J. T.; Kaczur, J. J.; Majsztrik, P. W.; Bocarsly, A. B. J. CO2 Util. 2014, 7, 1–5. Kapusta, S.; Hackerman, N. J. Electrochem. Soc. 1983, 130, 607–613. Innocent, B.; Liaigre, D.; Pasquier, D.; Ropital, F.; Léger, J.-M.; Kokoh, K. B. J. Appl. Electrochem. 2008, 39, 227–232. Innocent, B.; Pasquier, D.; Ropital, F.; Hahn, F.; Léger, J.-M.; Kokoh, K. B. Appl. Catal. B Environ. 2010, 94, 219–224. Komatsu, S.; Yanagihara, T.; Hiraga, Y.; Tanaka, M.; Kunugi, A. Denki Kagaku Oyobi Kogyo Butsuri Kagaku 1995, 63, 217–224. Zhang, H.; Ma, Y.; Quan, F.; Huang, J.; Jia, F.; Zhang, L. Electrochem. Commun. 2014, 46, 63–66. 26 ACS Paragon Plus Environment

Page 27 of 27

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis

(21) (22) (23) (24) (25) (26) (27) (28) (29) (30) (31) (32) (33) (34) (35) (36) (37) (38) (39) (40)

Chen, Y.; Li, C. W.; Kanan, M. W. J. Am. Chem. Soc. 2012, 134, 19969–19972. Li, C. W.; Ciston, J.; Kanan, M. W. Nature 2014, 508, 504–507. Min, X.; Kanan, M. W. J. Am. Chem. Soc. 2015, 137, 4701–4708. Li, C. W.; Kanan, M. W. J. Am. Chem. Soc. 2012, 134, 7231–7234. Ammar, I. A.; Darwish, S.; Khalil, M. W. Corrosion 1976, 32, 173–178. Burbank, J. J. Electrochem. Soc. 1957, 104, 693–701. Wang, M.; Liu, Y.; Zhang, D.; Yang, H. Electrochimica Acta 2012, 59, 441–448. Czerwiński, A.; Żelazowska, M.; Grdeń, M.; Kuc, K.; Milewski, J. D.; Nowacki, A.; Wójcik, G.; Kopczyk, M. J. Power Sources 2000, 85, 49–55. Visscher, W. J. Power Sources 1976, 1, 257–266. Sunderland, J. G. J. Electroanal. Chem. Interfacial Electrochem. 1976, 71, 341–345. Williams, D. E.; Wright, G. A. Electrochimica Acta 1976, 21, 1009–1019. Vivier, V.; Régis, A.; Sagon, G.; Nedelec, J.-Y.; Yu, L. T.; Cachet-Vivier, C. Electrochimica Acta 2001, 46, 907–914. Kim, K. S.; O’Leary, T. J.; Winograd, N. Anal. Chem. 1973, 45, 2214–2218. Dharmadhikari, V. S.; Sainkar, S. R.; Badrinarayan, S.; Goswami, A. J. Electron Spectrosc. Relat. Phenom. 1982, 25, 181–189. Pourbaix, M. Atlas of Electrochemical Equilibria; Burgers, W.g., Ed.Pergamon Press: London, 1974. Miller, F. A.; Wilkins, C. H. Anal. Chem. 1952, 24, 1253–1294. Taylor, P.; Sunder, S.; Lopata, V. J. Can. J. Chem. 1984, 62, 2863–2873. Packter, A. Krist. Tech. 1980, 15, 1237–1248. Conway, B. E.; Bockris, J. O. J. Chem. Phys. 1957, 26, 532–541. Quaino, P.; Juarez, F.; Santos, E.; Schmickler, W. Beilstein J. Nanotechnol. 2014, 5, 846– 854.

TABLE OF CONTENTS FIGURE


Probing the Mechanism of Aqueous CO2 Reduction on Post

Recommend Documents