Revisiting Electrochemical Reduction of CO2 on Cu Electrode: Where

Subscriber access provided by Kaohsiung Medical University

C: Surfaces, Interfaces, Porous Materials, and Catalysis 2

Revisiting Electrochemical Reduction of CO on Cu Electrode: Where Do We Stand about the Intermediates? Aamir Hassan Shah, Yanjie Wang, Abebe Reda Woldu, Lin Lin, Muzaffar Iqbal, David Cahen, and Tao He J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.8b05348 • Publication Date (Web): 23 Jul 2018 Downloaded from http://pubs.acs.org on July 26, 2018

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 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 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.

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 28 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

The Journal of Physical Chemistry

Revisiting Electrochemical Reduction of CO2 on Cu Electrode: Where Do We Stand about the Intermediates? Aamir Hassan Shah,†,‡ Yanjie Wang,†,‡ Abebe Reda Woldu,†,‡ Lin Lin,† Muzaffar Iqbal,† David Cahen,§ and Tao He*,†,‡ †

CAS Laboratory of Nanosystem and Hierarchical Fabrication, CAS Center for Excellence in Nanoscience, National Center for Nanoscience and Technology, Beijing 100190, China ‡

University of Chinese Academy of Sciences, Beijing 100049, China

§

Department of Materials & Interfaces, Weizmann Institute of Science, Israel

*To whom correspondence should be addressed; Tel: +86-10-82545655; Fax: +86-10-62656765 E-mail: [email protected]

ABSTRACT Electrochemical CO2 reduction on Cu electrode has attracted attention of many researchers in the last decades, because of its potential to generate significant amounts of hydrocarbons at high reaction rates over sustained periods of time. As a result, substantial effort has been devoted to determine the unique catalytic performance of Cu and to elucidate the mechanism through which hydrocarbons are formed. Here we report new insights into CO2 reduction on Cu by electrochemical impedance spectroscopy (EIS) in terms of adsorption/desorption of the reduction intermediates. The potential dependence of charge transfer kinetics is discussed on the basis of EIS results. We revisit the mechanism of the formation of hydrocarbons, taking into account the pH adjacent to the electrode surface, adsorption of HCO3‒ and CO32‒, and the role of active hydrogen. In addition to the enol-like intermediate, proposed previously, we proposed that *COOH• radicals,

1

ACS Paragon Plus Environment

The Journal of Physical Chemistry 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 28

originating from the active involvement of HCO3‒ and/or CO32‒ upon reduction are key intermediates for the formation of a variety of C2 and C3 products. Thus, our results provide an additional crucial guideline for the design of future catalysts that can efficiently and selectively reduce CO2 into value-added chemicals.

1.

INTRODUCTION

The continuously increasing amount of carbon dioxide (CO2) in the earth’s atmosphere is an issue of global concern, as CO2 is considered one of the major contributors to the greenhouse effect.1 Electrocatalytic CO2 reduction can help to achieve a sustainable redox cycle for energy storage and conversion to close the anthropogenic carbon cycle, but to that end efficient CO2 reduction catalysts are required to convert CO2 into fuels and water. Electrochemical CO2 reduction has been intensively studied in the last decades.2-5 Developing efficient CO2 reduction electrocatalysts is an important part of the quest to realize direct electrolytic fuel synthesis using a renewable electricity source.4 Scientists over last three decades have identified materials that are capable of electrochemically reducing CO2 in aqueous solution.6 Despite the great potential and significant efforts, truly efficient electrocatalytic reduction of CO2 remains elusive. Ever since Hori made the milestone discovery in 1985 that copper (Cu) has unique ability to reduce CO2 to hydrocarbons with Faradaic efficiencies that are better than those of other electrocatalysts, Cu has been one of the most studied material for catalyzing CO2 reduction.7 Thus, understanding its unique ability to form hydrocarbons may help design new catalysts that are both active at low overpotentials with high efficiency, and have better product selectivity than existing ones. Substantial efforts have been devoted to understand the special reactivity of Cu for CO2 reduction and a number of reports summarize the literature associated with CO2 reduction on Cu surfaces.8-

2


Page 3 of 28 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


10

Earlier work reported CO as a key intermediate in the formation of hydrocarbons, which is now

widely accepted in the literature.11 However, the complete reduction mechanism is still unclear as illustrated by the observation of 16 different products formed by CO2 reduction.12 The products include a broad mixture of aldehydes, ketones, carboxylic acids and alcohols, showing the complexity of this reaction. It is understood that quantity and selectivity of the reduction products rely on many factors, such as solvent, nature and pH of electrolyte, electrode material, and applied bias.13-14 Theoretical studies have indicated that the optimum bulk pH for CO2 reduction is close to 7. A higher or lower local pH value results in higher Nernstian loss, resulting in decrease of the actual bias that is applied to drive the CO2 reduction kinetics at the cathode and water oxidation at the anode.15 The distribution between molecular CO2 and other forms of oxidized carbon like carbonic acid (H2CO3), bicarbonates (HCO3‒) and carbonates (CO32‒), is strongly pH dependent. A higher local pH leads to a decrease in CO2 concentration, because of its conversion into HCO3‒ and CO32‒ via reaction with the hydroxyl anions; while the formation of HCO3‒ and CO32‒ occurs at much higher rate than that of CO2 reduction. The most accepted theory for CO2 reduction is formation of the carbon dioxide radical anion intermediate (CO2•‒), which is considered the ratelimiting step for CO2 reduction.16-17 However, there is still controversy about how the CO2•‒ is formed. Recently, Narayanan et al. claimed the absence of CO2 in the aqueous solution system,18 while others believe that CO2 molecules can exist in the aqueous system as H2CO3 decomposes reversibly with a life time of 60 ms to CO2 and water in aqueous solution at room temperature with a first order rate constant.19-21 However, in a near-neutral electrolyte the pH near the cathode increases under more negative potential during CO2 reduction, which results in a low [CO2] near the cathode. For example, 0.1 M aqueous solution of CO2-saturated potassium bicarbonate (KHCO3) has a bulk pH of 6.8, while the pH increases to 9.5 near the cathode surface under an

3



Page 4 of 28

applied potential of –1.15 V versus RHE.15, 22 M. Dunwell and coworkers recently proposed the role of bicarbonates in CO2 reduction on Au electrode surface23 — previously considered only as a buffering agent or proton source14, 24 — that bicarbonate, through equilibrium exchange with dissolved CO2, rather than the supplied CO2, is the primary source of carbon in the CO formed at the Au electrode. However, it is still unclear what is the actual carbon source for the formation of the CO2•‒ radical on Cu electrode surface and, specifically, which intermediate plays a key role in electrocatalytic CO2 reduction. It is important to know the fundamental behavior of the electrode/electrolyte interface before any technical consideration on the implementation of an electrolytic system for CO2 reduction to products is made. The primary step towards practical performance is definitely the formulation of the mechanism that can explain the observed product distribution and onset potentials, while taking into account the effect of pH on and near the cathode surface. Albeit several reports are available that considered the local pH effect on product formation and distribution, surprisingly, no reports have considered the local pH effect and proposed a detailed mechanism for CO2 reduction on Cu electrode in aqueous solution.8, 12-13 Cyclic voltammetry and related electrochemical techniques are often-used tools to study electron transfer reactions, because the method can indeed provide valuable insights into redox reaction mechanisms, energetics and rates and, probably also because of the ease of implementation, low cost of the system and its operation. Whatever an electroanalytical system or method is used for investigation, electrochemical impedance spectroscopy (EIS) has its own strengths as it can work in the conditions of both steady state and those far from the steady state. It is of utmost importance since it can show perturbations in large potential window, especially for electrochemical reactions. EIS is also one of the most powerful and versatile nondestructive tools capable to analyze the 4


Page 5 of 28 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


complex electrochemical process over a wide range of frequencies. Moreover, the kinetic data of redox reaction, behavior of electrode/electrolyte interface, adsorption/desorption behavior of intermediates on the electrode surface, and nature of redox process can be easily determined using EIS technique. Herein we performed detailed EIS analysis to determine the adsorption/desorption behavior of CO2, HCO3‒ and CO32‒ on the Cu electrode surface. Depending on the pH, three different systems were established, i.e., CO2 (H2CO3 ↔ CO2 + H2O), HCO3‒ and CO32‒. The potential at which CO2 reduction starts was observed, and the potential at which all the three systems behave similarly was determined too. Based on our experimental results we proposed a mechanism for reduction of the CO2-saturated 0.1 M KHCO3 electrolyte system with bulk pH of 6.8, which accounts for the effect of pH on the electrode surface, potential-dependent product distribution, C-C and C-C-C coupling, intermediates and rate-determining steps of different CO2 reduction processes.

2.

Experimental

2.1.

Materials

Potassium bicarbonate (99.999% metal basis) and potassium hydroxide (99.999% metal basis) were purchased from Sigma-Aldrich. Copper foil (99.9999% metal basis) was bought from Alfa Aesar. Carbon dioxide (99.999%) and argon (99.999%) were purchased from Beiwen. Electrolyte solutions were prepared with 18.2 MΩ deionized water obtained from a Millipore system. 2.2.

Electrode and electrolyte preparation

Copper foil was cut into electrodes with 1×1 cm2 square and then cleaned by sonicating for 30 min in acetone, followed by isopropyl alcohol and finally in deionized water. It was then electropolished in 85% phosphoric acid under a bias of 2 V for 30 min, followed by rinsing with 5



Page 6 of 28

deionized water and drying with a stream of argon. To prepare carbonic acid solution, 0.1 M potassium bicarbonate (KHCO3) was purged with argon for 30 min, followed by purging by CO2 for 60 min to maintain pH of 6.8. The bicarbonate solution was prepared by purging 0.1 M KHCO3 with argon for 30 min (pH 9.0), and the carbonate solution was prepared by adding the potassium hydroxide in 0.1 M KHCO3 to maintain the pH of 10.8. 2.3.

Electrochemical measurements

Electrochemical measurements were carried out using Zahner IM6 potentiostat (Zahner Elektrik, Kronach). The linear scan voltammetry and electrochemical impedance spectroscopy (EIS) were performed in homemade electrochemical cell. Before any electrochemical measurement Cu electrode was electrochemically polarized at a slight negative potential to avoid any ambiguity related to oxide formation on the electrode surface. Briefly, the working electrode was parallel to counter electrode to ensure uniform potential distribution across the electrode surface. The EIS spectra were collected in the frequency range of 0.1 MHz to 1.0 mHz with an AC amplitude of 100 mV under different applied bias. All the above mentioned experiments were repeated at least three times so as to make sure the results are reasonable and reproducible. A Z-view software was used to fit the Nyquist plots. For all of the experiments, the platinum foil was used as counter electrode and Ag/AgCl (Innovative Instruments, Inc.) was used as reference electrode. The data were converted to RHE scale by using the following equation. 𝐸(𝑣𝑠 𝑅𝐻𝐸) = 0.197 + 0.059(𝑝𝐻) + 𝐸(𝑣𝑠 𝐴𝑔/𝐴𝑔𝐶𝑙)

3.

RESULTS AND DISCUSSION

3.1.

Voltammetric Characteristics

6


Page 7 of 28

Figure 1 shows linear scanning voltammetric (LSV) results of Cu foil in H2O, CO2, HCO3‒ and CO32‒-containing electrolytes, at a 10 mV/sec scan rate. For the Ar-purged H2O system it is found that there is almost no current flow in the given potential window. The reduction current for CO2saturated 0.1 M KHCO3 (pH 6.8) starts at ‒0.35 V (referenced to RHE, unless stated otherwise) with a shoulder at ‒0.63 V, and then increases gradually. As reported previously, the adsorption of CO on a Cu electrode occurs at ‒0.6 V and CO starts to desorb at ‒0.8 V;25 hence, the shoulder at ‒0.63 V is assigned to partial adsorption of CO on the electrode surface due to CO2 reduction and the gradual increase in the shoulder is attributed to conversion of adsorbed CO into hydrocarbons (see mechanism section) and desorption from the electrode surface. 0.0 -0.2

J /mA.cm-2

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


-0.4 -0.6 -0.8

Ar Purged Water (pH 7.0)

-1.0

CO2 saturated KHCO3 (pH 6.8) Ar Purged HCO3- (pH 9.2)

-1.2

Ar Purged CO32- (pH 10.8)

-1.4 -0.8

-0.7

-0.6

-0.5

-0.4

-0.3

-0.2

-0.1

E /V versus RHE Figure 1. Linear scan voltammograms (LSV) of H2O, CO2 saturated 0.1 M KHCO3, Ar purged 0.1 M KHCO3 and 0.1 M K2CO3 at scan rate of 10 mV/sec

With HCO3‒ and CO32‒ solutions, reduction current starts at about ‒0.3 V and then increases sharply. The more negative onset observed in CO2-saturated 0.1 M KHCO3 system is ascribed to the interference caused by some intermediate species formed in case of CO2 reduction. The less

7



Page 8 of 28

negative onset in HCO3‒ and CO32‒ indicates that the reduction of HCO3‒ and CO32‒ may have different intermediate formation from that for CO2 reduction, as will be discussed below. Similarly, absence of any shoulder in case of HCO3‒ and CO32‒ confirms that CO is not the reduction product of HCO3‒ and CO32‒ reduction. This can be further supported by the formation of potentialdependent products, available in previous reports (Figure S1).10, 12 3.2.

Electrochemical Impedance Spectroscopy Characteristics

The Nyquist plots of CO2-saturated 0.1 M KHCO3 in Figure 2 show potential-dependent charge transfer behavior on the electrode surface. The positive and less negative potentials (i.e., in the range of 0.04 to ‒0.28V) were applied to study pre-adsorption of electroactive species on the electrode surface. The hydrated cations can act as a pH buffer near the cathode and have an impact on stabilizing the adsorption of CO2, HCO3‒ and CO32‒ at the cathode. The hydrated potassium cations (K+) are stable against hydrolysis in the bulk of the electrolyte, but can undergo hydrolysis in the vicinity of a cathode as a result of Columbic interactions with the negatively charged cathode.22 The equivalent circuits used for data fitting are shown in Figure S2, in which charge transfer resistance (RCT) refers to the barrier that the electrons must cross to reach the electroactive species adsorbed on the electrode surface. Figure 2a shows the two semicircles are well separated at a bias of 0.04 V, while decreasing the bias leads to merging of the two semicircles into one. The second semicircle completely merges into the first one at ‒0.28 V bias. The decrease in the semicircle radius at more negative potentials indicates facile charge transfer kinetics under such conditions. These observations imply that the charge transfer starts taking place even at a relatively less negative potential, while the activation barrier is still significant enough to give almost no detectable products.

8


Page 9 of 28

6.0

1.5

(a) 0.04 V -0.08 V -0.13 V -0.19 V -0.25 V -0.28 V

4.5 0.01

3.0 0.00

0.015

0.030

0.045

(b) -0.31 V -0.35 V -0.40 V -0.48 V -0.52 V -0.57 V -0.61 V

0.02

1mHz

0.1MHz

0.04

1.2

-Z(Im)/ k

0.02

-Z(Im)/ k

1.5

0.0

0.9 0.00

0.025

0.050

0.075

0.6

0.1MHz

1mHz

0.3 0.0

0.0

1.5

3.0

4.5

6.0

Z(Re) / k

7.5

0.0

9.0

0.5

1.0

1.5

2.0

Z(Re) / k

0.10

(c) -0.66 V CO2/HCO-3

0.08

-0.66 V Bicarbonates -0.66 V Carbonates

-Z(Im)/ k

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


0.06 0.04

1mHz

0.1MHz 0.02 0.00 0.03

0.06

0.09

0.12

0.15

Z(Re) / k Figure 2. EIS Nyquist plots of (a) CO2 saturated 0.1 M KHCO3 electrolyte (pH 6.8) at and below ‒0.28 V (b) CO2 saturated 0.1 M KHCO3 electrolyte (pH 6.8) at and below ‒0.61 V and (c) comparison between CO2, HCO3‒ and CO32‒ systems at ‒0.66 V. Green lines are the fitted spectra, and insets show the zoom-in portion at higher frequencies of corresponding plots

Figure 2b displays the Nyquist plots at bias voltages between ‒0.31 to ‒0.61 V. The sudden decrease in the RCT and appearance of a new semicircle at ‒0.31 V marks this bias as a critical potential at which CO2 reduction intermediate(s) start(s) to be formed on the electrode surface. A 9



new semicircle starts at ‒0.31 V, and becomes more pronounced as the potential becomes more negative. In addition, the adsorption phenomenon on the electrode surface becomes evident from the negative values of the imaginary component of the impedance (Z″) at low frequencies, which is because of the adsorption of intermediates arising from the reduction of CO2, HCO3‒ and/or CO32‒. The significant value of Z″ before the appearance of new semicircle at ‒0.31 V and more negative potentials indicates that faradaic adsorption of intermediates occurs on the electrode surface. EIS at some selected potentials was also performed in aqueous solutions with different pH values so as to compare the EIS characteristics for the systems that contain mainly HCO3‒ or mainly CO32‒. It is found that the charge transfer behavior is completely different at less negative bias (Figure S3) from that for the CO2-saturated 0.1 M KHCO3 system, while all the three systems behave similarly at a more negative bias like ‒0.66 V (Figure 2c). These results indicate the existence of only HCO3‒ and/or CO32‒ species on the electrode surface at more negative potentials, implying the local pH value adjacent to electrode surface is different from the bulk one under more negative bias.

60

75

(a)

30 15 0

-3.0

45

-0.31 V -0.35 V -0.40 V -0.44 V -0.48 V -0.52 V -0.57 V -0.61 V

4 0

o

o

45

(b)

8

-Phase angle/

0.04 V -0.08 V -0.13 V -0.19 V -0.25 V -0.28 V

60

-Phase angle/

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 28

-3

-2

-1

30

15

0 -1.5

0.0

1.5

3.0

4.5

-3.0

-1.5

0.0

1.5

3.0

4.5

Log Frequency

Log Frequency

10


Page 11 of 28

35

(c)

-0.66 V CO2/HCO3o

28

-Phase angle/

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


-0.66 V Bicarbonates -0.66 V Carbonates

21

14

7

0 -3.0

-1.5

0.0

1.5

3.0

4.5

Log Frequency Figure 3. EIS Bode phase angle plots of (a) CO2-saturated 0.1 M KHCO3 electrolyte (pH 6.8) at ‒0.28 V and smaller bias voltages; (b) CO2-saturated 0.1 M KHCO3 electrolyte (pH 6.8) at ‒0.61 V and smaller bias voltages (inset shows zoom-in portion of corresponding spectra at lower frequencies); and (c) comparison between CO2, HCO3‒ and CO32‒ systems at ‒0.66 V.

To interpret the phenomena observed in Figure 2, the plots of phase angle and impedance modulus versus log frequency are shown in Figure 3 and Figure 4, respectively. Figure 3a corresponds to the phase angle plots observed at potential of and less negative than ‒0.28 V. The two clearly distinct peaks at 0.04 V bias at central and low frequency correspond to the two semicircles observed in the Nyquist plot (Figure 2a). While increasing the negative potential the small peak at low frequencies starts merging into the large peak observed at central frequencies. The negative values of phase angles at low frequencies in the potential range of ‒0.13 to ‒0.44 V are due to the adsorption of intermediates formed during the reduction process. Moreover, the appearance of a new small peak at low frequencies starts at ‒0.31 V, and becomes more distinguishable at more negative potentials (Figure 3b), which corresponds to charge transfer to the intermediates adsorbed faradaically on the electrode surface. Figure 3c shows the comparison of phase angle plots of 11



solutions that contain mainly CO2, HCO3‒ or CO32‒. The CO2-saturated 0.1 M KHCO3 system shows a shoulder in the high frequency region, which corresponds to the CO formation and, thus, complements the above LSV results, while the second and third peak at intermediate and lower frequency can be correlated with HCO3‒ and CO32‒ systems. These results imply that the local pH value increases during the EIS measurements, leading to depletion of CO2 and accumulation of HCO3‒ and/or CO32‒ on the electrode surface.

0.04 V -0.08 V -0.13 V -0.19 V -0.25 V -0.28 V -0.31 V -0.35 V -0.4 V -0.44 V -0.48 V -0.52 V -0.57 V -0.61 V -0.66 V

2

8

1

6

[Z]/k

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 28

0 -2

0

2

4

2

0 -2

-1

0

1

2

3

Log Frequency Figure 4. EIS Bode impedance modulus plot of CO2 saturated 0.1 M KHCO3 electrolyte (pH 6.8) at different applied bias (inset is zoom-in on lower left part of the main figure).

It can be seen from Figure 4 that the impedance modulus Bode diagrams show almost the same value at high frequencies at different potentials, corresponding to θ = 0 in the phase angle Bode plot, which is the electrolyte resistance between working and reference electrodes. The system shows two slopes (i.e., two time constants) upon decreasing frequency, associated with a kinetic process (high frequencies) and a diffusion process (low frequencies). The decrease in impedance indicates facile charge transfer kinetics at more negative potentials.

12


Page 13 of 28 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


3.3.

Surface Coverage of Faradaic Adsorption

The extent of faradaic adsorption on an electrode surface can also be determined by EIS. It is wellknown that several factors can have an impact on precise measurement of interfacial capacitance, such as micro-roughness of the electrode and adsorption/desorption of the electrolyte species. Accordingly, the imaginary component of impedance (Z″) is assumed to represent the dynamic process occurring at the interface induced by the adsorption.26 Here the relationship used to determine the surface coverage percentage is established by making three assumptions, the value of Z″ on bare Cu electrode should be zero, the maximum value of Z″ (denoted as Z″max) is observed when electrode surface is completely covered by the analyte, and the value of Z″ observed before the appearance of a new semicircle at ‒0.31 V (Figure 2b) and more negative potential (denoted as Z″ads) is due to faradaic adsorption of reactant species on the electrode surface. The surface coverage percentage for the species adsorbed on the electrode surface is determined by the ratio of Z″ads to Z″max. The obtained surface coverage with respect to potential shows that almost 20~30% of the electrode surface is covered as a result of faradaic adsorption during reduction (Figure 5). Again, no faradaic adsorption occurs under bias voltage that is less negative than –0.3 V. We noted earlier that H2CO3 and HCO3‒ are the dominant species at the near neutral pH range,22 which is generally employed for electrochemical reduction of CO2. The molar fraction of H2CO3 becomes less significant in alkaline media, whereas HCO3‒ and/or CO32‒ become the major species.27 In this case, the reactant that is mainly converted into carbonaceous products is under debate, although there is some evidence that HCO3‒ is the (dominant) reactant.28-31 On the basis of the EIS results (Figure 2C) and in terms of a recent report about pH determination on the electrode surface,22 the previously proposed mechanism is revisited and, specifically, the potential dependent charge-transfer mechanism and earlier reported product distribution is reanalyzed.12 13



Obviously, our potential-dependent EIS results, shown in Figure 2, indicate facile charge transfer kinetics under more negative bias and, thereby, formation of various reduction products is not surprising.

32 30

Surface Coverage / %

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 28

28 26 24 22 -0.3

-0.4 -0.5 -0.6 Potential /V versus RHE

-0.7

Figure 5. Potential-dependent surface coverage derived from data shown in Figure 2b.

The first considered major product is hydrogen (H2). H2 evolution appeared at ‒0.6 V decreases with increasingly negative potential (Figure S1). This decrease in the current efficiency of H2 is attributed partially to less availability of protons on the electrode surface due to the high local pH value at more negative potentials, and partially to the availability of adsorbed H atoms in that they can form other major hydrocarbons products of methane, ethylene and formate. Such decrease continues until ‒1.05 V. The H2 formation starts increasing again by further increasing (making more negative) the potential, which is ascribed to the decrease in the formation of C2 and C3 hydrocarbon products at more negative potentials and, hence, more adsorbed H atoms can participate in the H2 evolution.

14


Page 15 of 28 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


Two other major products of CO2 reduction under less negative bias (such as ‒0.60 to ‒0.90 V) are formate (HCOO‒) and carbon monoxide (CO). The HCOO‒ starts increasing with increase of the potential and reaches the maximum at around ‒0.90 V; while the amount of CO decreases continuously with applying more negative potentials, which is partly attributed to the further reduction of CO to hydrocarbons (see mechanism section), and partly because HCO3‒ and CO32‒ start dominating on the electrode surface, since the pH increases due to the reduction reaction (Figure S4).27 The decrease in both HCOO‒ and CO observed at even more negative potentials can be assigned to decrease in the amount of CO2 molecules on the electrode surface as a result of mass transfer limitation and increase in local pH at more negative potentials.15, 22 Such decrease continues by increasing the potential, with almost complete absence of these products under a bias at and more negative than about ‒1.05 V. Methane (CH4) and ethylene (C2H4) are also found as major products of CO2 reduction, while their amount under less negative bias (such as ‒0.60 to ‒0.90 V) are less than those of HCOO‒, CO, and H2. Moreover, the lower amount of CH4 as compared to C2H4 at less negative potentials (such as ‒0.60 to ‒0.90 V) is attributed to different kinetics of CH4 and C2H4 formation and different ratedetermining steps leading to each of these products (see mechanism discussion).8 The amount of CH4 increases gradually when the potential moves to more negative values than ‒0.95 V; whereas the amount of C2H4 first increases steadily when the potential becomes more negative because dimerization is more facile than charge transfer and then starts to decrease at around ‒1.05 V due to the decrease in C˗C coupling. This reversal may be because the surface bound intermediates are more likely to be reduced to CH4 and/or desorbed at more negative potentials, as manifested by the decreased surface coverage shown in Figure 5.

15



Page 16 of 28

The observation of other intermediate and minor products like alcohols, aldehydes and ketones at ‒0.8 V and more negative potentials are considered the reduction products of HCO3‒ and/or CO32‒ because at more negative potentials the significant CO2 molecules are absent on the electrode surface due to the high local pH.22 Further evidence for the absence of CO2 molecules under a bias more negative than ‒1.0 V is that almost no CO is produced under such circumstance. In addition, direct reduction of HCO3‒ on a Cu surface has been proven experimentally,27 indicating the favorable reduction of HCO3‒ rather than its decomposition into CO2. It is reported previously that alkaline electrolyte can facilitate the formation of C2 over C1 products on different Cu surface structures.32-33 Based on the aforementioned EIS results, we suggest that the reduction of HCO3‒ can occur on the Cu electrode because of the alkaline environment induced by more negative bias. The amount of intermediate and minor products increases with increasing negative potential mainly due to less charge transfer resistance at these bias voltages; further increase in the potential leads to decrease in C˗C coupling and, hence, a sudden drop in the formation of these products. 3.4.

Proposed Redox mechanism

Based on the above discussions, we revisit the redox mechanism of all the reported products. The one for C1 species is put forward with some modifications, based on that proposed previously.13, 34

The reduction mechanism of CO2 into formate and CO is schematically shown in Figure 6. The

first intermediate for ethylene (C2H4) and all the C1 products (except methanol) is *CO2•‒, which can be further reduced via two pathways, either by reduction of one electron with one proton to produce formate, or via protonation to generate the *COOH intermediate that can be further reduced by one electron with one proton to form CO. The *COOH intermediate formation is also possible through a concerted proton-electron transfer to CO2 as shown in Figure 6.13

16


Page 17 of 28 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


Figure 6. Proposed mechanisms of formation of some products of CO2 reduction: formate, CO, CH4 and C2H4.

The pH-independent *CO dimerization has been proposed previously, i.e., two *CO molecules dimerize, accompanied with one-electron reduction to form the *C2O2‒ anion35, followed by subsequent reduction of *C2O2‒ into C2H4.36 It has been demonstrated theoretically that *CO dimerization is kinetically unfavorable owing to a kinetic barrier and, thus, *CO protonation is critical to make it exergonic to achieve favorable kinetics.37 The barrier to *CO dimerization is well beyond the range required for favorable C-C coupling kinetics.37 Hence, the formation of the *CH2OH intermediate via reduction of adsorbed CO by three electrons with three protons cannot be ruled out. The subsequent reduction of *CH2OH intermediate by one electron with one proton can give rise to the formation of *CH2• radical (Figure 6). As the similar onset potentials of CH4 and C2H4 (Figure S1) may indicate the reduction of a common intermediate, the dimerization of the *CH2• radical can result in the formation of C2H4,38-39 whereas its further reduction by two

17



Page 18 of 28

electrons with two protons can lead to the production of CH4. The effective barrier for C2H4 formation from CH2 dimerization is reported as 0.21 eV at 0 V (RHE), and the barrier for CH2 reduction to CH3 is reported to be 0.55 eV at 0 V (RHE) and 0 eV at ‒1.15 V (RHE).39 Albeit the rate constant favors C2H4 production under less negative bias, the relative surface coverage of CH2 and adsorbed H should also have an impact on the selectivity and, thereby, may promote CH4 formation under more negative bias. According to the above EIS results (Figure 2c) and previous reports,15, 22 the predominant species on a Cu electrode surface under more negative bias are HCO3‒ and/or CO32‒ because of the high pH value adjacent to the electrode surface. Since all of the C2 (except C2H4) and C3 species, as well as CH3OH, start forming under high bias, the starting reduction material is considered HCO3‒ and/or CO32‒, as presented in Figures 7 and 8. The one-electron reduction of the HCO3‒ and/or CO32‒ results in the formation of the *CO2•‒ radical, which can be further protonated to produce the *COOH• radical. It is suggested that the *COOH• radical is the key starting intermediate, which can lead to the formation of CH3OH via reduction by five electrons with five protons under more negative bias (Figure 7). An enol-like intermediate has been detected previously in solution and is thought to play a key role in the formation of all C2 and C3 products, even though its origin is unclear.12 Here it is suggested that the *COOH• radical can dimerize to produce the *(COOH)2 intermediate (Figure 7), which can be reduced via two electrons with two protons to produce the *(C(OH)2)2 intermediate, followed by further reduction via four electrons with four protons to generate glycolaldehyde (HOCH2-CHO), or by reduction through two electrons with two protons to form *(CH(OH)2)2 that can eventually lead to the formation of glyoxal (OCHCHO). In addition, the reduction of *(C(OH)2)2 via four electrons with four protons can also give rise to the production 18


Page 19 of 28 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


of *(HCOH)2 intermediate, which can result in the generation of ethylene glycol (CH2OH)2 via reduction by two electrons with two protons or ethanol (C2H5OH) through reduction by three electrons with three protons. The reduction of *(C(OH)2)2 via one electron with one proton can result in the formation of the *HCOHCOOH intermediate, which can produce acetaldehyde (CH3CHO) via reduction by five electrons with five protons or acetic acid through reduction by three electrons with three protons. In addition, the adsorbed enol-like intermediate can react with *COOH• radical to form the C3 intermediate like *C(OH)2COC(OH)2 via reduction by one electron with one proton, or produce the *C(OH)2COHC(OH)2 intermediate via reduction by two electrons with two protons (Figure 8). The former can give rise to the formation of acetone (CH3COCH3) via reduction by ten electrons with ten protons, or hydroxyacetone (C3H6O2) via reduction by eight electrons with eight protons; the latter can lead to the generation of 1-propanol (C3H8O) via reduction by eleven electrons with eleven protons, or propanal or alkyl alcohol via reduction by nine electrons with nine protons.

4.

Conclusion

In summary, we have performed detailed voltammetric and EIS analysis on CO2, HCO3‒ and CO32‒ systems in aqueous electrolyte. The critical bias (0.13V) at which the adsorption of electroactive species starts and critical bias (0.31 V) for the onset of charge transfer to the adsorbed species are determined. The potential (0.61V) at which all the three CO2, HCO3‒ and CO32‒ systems behave similarly on the electrode surface is also determined. The percent surface coverage of the adsorbed species on electrode surface is evaluated.

19



Page 20 of 28

Figure 7. Proposed formation mechanism of C2 species and methanol upon CO2 reduction.

20


Page 21 of 28 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


Figure 8. Proposed formation mechanism of C3 species upon CO2 reduction.

Taking into account for the first time the effect of pH adjacent to electrode surface, we put forward a detailed mechanism that accounts for the effect of pH on the electrode surface, potentialdependent product distribution, C-C and C-C-C coupling, intermediates and rate-determining steps of different reduction processes. The *COOH• radical intermediate is considered a key intermediate for the formation of various C2 and C3 products. Our results suggest that the mechanism by which C-C and C-C-C bonds are formed is dependent on the applied bias and pH 21



Page 22 of 28

value adjacent to the electrode surface. Our work should lead to clearer mechanistic insight, to help future catalyst searching.

Associated Content: Supporting Information. All supporting information are available free of charge via the Internet at http://pubs.acs.org, including current efficiency for the products of CO2 reduction on Cu electrode in 0.1 M KHCO3, as a function of potential, equivalent circuit used for fitting the EIS data, EIS Nyquist plots of Arpurged 0.1 M KHCO3 electrolyte with pH values, and species fraction of CO2, HCO3‒ and CO32‒ at different pH values in aqueous system. AUTHOR INFORMATION Corresponding Author * Tao He, E-mail: [email protected] ORCID David Cahen: 0000-0001-8118-5446. Tao He: 0000-0001-6336-2402. Notes The authors declare no competing financial interest.

Acknowledgements:

22


Page 23 of 28 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


This work was supported by the Ministry of Science and Technology of China (2015DFG62610) and The Belt and Road Initiative by Chinese Academy of Sciences. At the Weizmann Institute of Science this research was supported in part by the China - Israel program of the Israel Ministry of Science, Technology and Space. A. H. Shah thanks The CAS-TWAS President’s Fellowship for International PhD Students. References: 1.

Figueiredo, M. C.; Ledezma-Yanez, I.; Koper, M. T. In Situ Spectroscopic Study of CO2

Electroreduction at Copper Electrodes in Acetonitrile. ACS Catal. 2016, 6, 2382-2392. 2.

Qiao, J.; Liu, Y.; Hong, F.; Zhang, J. A Review of Catalysts for the Electroreduction of

Carbon Dioxide to Produce Low-Carbon Fuels. Chem. Soc. Rev. 2014, 43, 631-675. 3.

Kuhl, K. P.; Hatsukade, T.; Cave, E. R.; Abram, D. N.; Kibsgaard, J.; Jaramillo, T. F.

Electrocatalytic Conversion of Carbon Dioxide to Methane and Methanol on Transition Metal Surfaces. J. Am. Chem. Soc. 2014, 136, 14107-14113. 4.

Li, C. W.; Kanan, M. W. CO2 Reduction at Low Overpotential on Cu Electrodes Resulting

from the Reduction of Thick Cu2O Films. J. Am. Chem. Soc. 2012, 134, 7231-7234. 5.

Hori, Y.; Murata, A.; Yoshinami, Y. Adsorption of CO, Intermediately Formed in

Electrochemical Reduction of CO2, at a Copper Electrode. J. Chem. Soc. Faraday Trans. 1991, 87, 125-128. 6.

Chen, Y.; Kanan, M. W. Tin Oxide Dependence of the CO2 Reduction Efficiency on Tin

Electrodes and Enhanced Activity for Tin/Tin Oxide Thin-Film Catalysts. J. Am. Chem. Soc. 2012, 134, 1986-1989.

23



7.

Page 24 of 28

Hori, Y.; Kikuchi, K.; Suzuki, S. Production of CO2 and CH4 in Electrochemical Reduction

of CO2 at Metal Electrodes in Aqueous Hydrogencarbonate Solution. Chem. Lett. 1985, 14, 16951698. 8.

Gattrell, M.; Gupta, N.; Co, A. A Review of the Aqueous Electrochemical Reduction of

CO2 to Hydrocarbons at Copper. J. Electroanal. Chem. 2006, 594, 1-19. 9.

Whipple, D. T.; Kenis, P. J. A. Prospects of CO2 Utilization Via Direct Heterogeneous

Electrochemical Reduction. J. Phys. Chem. Lett. 2010, 1, 3451-3458. 10.

Hori, Y.; Murata, A.; Takahashi, R. Formation of Hydrocarbons in the Electrochemical

Reduction of Carbon Dioxide at a Copper Electrode in Aqueous Solution. J. Chem. Soc. Faraday Trans. 1989, 85, 2309-2326. 11.

Hori, Y.; Murata, A.; Takahashi, R.; Suzuki, S. Electroreduction of Carbon Monoxide to

Methane and Ethylene at a Copper Electrode in Aqueous Solutions at Ambient Temperature and Pressure. J. Am. Chem. Soc. 1987, 109, 5022-5023. 12.

Kuhl, K. P.; Cave, E. R.; Abram, D. N.; Jaramillo, T. F. New Insights into the

Electrochemical Reduction of Carbon Dioxide on Metallic Copper Surfaces. Energy Environ. Sci. 2012, 5, 7050-7059. 13.

Kortlever, R.; Shen, J.; Schouten, K. J. P.; Calle-Vallejo, F.; Koper, M. T. Catalysts and

Reaction Pathways for the Electrochemical Reduction of Carbon Dioxide. J. Phys. Chem. Lett. 2015, 6, 4073-4082. 14.

Hori, Y. Electrochemical CO2 Reduction on Metal Electrodes. In Modern Aspects of

Electrochemistry, Springer NewYork; NewYork., 2008; pp 89-189.

24


Page 25 of 28 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


15.

Singh, M. R.; Clark, E. L.; Bell, A. T. Effects of Electrolyte, Catalyst, and Membrane

Composition and Operating Conditions on the Performance of Solar-Driven Electrochemical Reduction of Carbon Dioxide. Phys. Chem. Chem. Phys. 2015, 17, 18924-18936. 16.

Chiesa, M.; Giamello, E.; Che, M. Epr Characterization and Reactivity of Surface-

Localized Inorganic Radicals and Radical Ions. Chem. Rev. 2010, 110, 1320-1347. 17.

Preda, G.; Pacchioni, G.; Chiesa, M.; Giamello, E. Formation of CO2•− Radical Anions

from CO2 Adsorption on an Electron-Rich Mgo Surface: A Combined Ab Initio and Pulse Epr Study. J. Phys. Chem. C. 2008, 112, 19568-19576. 18.

Narayanan, H.; Nair, M.; Viswanathan, B. On the Current Status of the Mechanistic

Aspects of Photocatalytic Reduction of Carbon Dioxide. Ind. J. Chem. 2017, 56, 152-169. 19.

Mills, G. A.; Urey, H. C. The Kinetics of Isotopic Exchange between Carbon Dioxide,

Bicarbonate Ion, Carbonate Ion and Water1. J. Am. Chem. Soc. 1940, 62, 1019-1026. 20.

Roughton, F. The Kinetics and Rapid Thermochemistry of Carbonic Acid. J. Am. Chem.

Soc. 1941, 63, 2930-2934. 21.

Pines, D.; Ditkovich, J.; Mukra, T.; Miller, Y.; Kiefer, P. M.; Daschakraborty, S.; Hynes,

J. T.; Pines, E. How Acidic Is Carbonic Acid? J. Phy. Chem. B. 2016, 120, 2440-2451. 22.

Singh, M. R.; Kwon, Y.; Lum, Y.; Ager III, J. W.; Bell, A. T. Hydrolysis of Electrolyte

Cations Enhances the Electrochemical Reduction of CO2 over Ag and Cu. J. Am. Chem. Soc. 2016, 138, 13006-13012. 23.

Dunwell, M.; Lu, Q.; Heyes, J. M.; Rosen, J.; Chen, J. G.; Yan, Y.; Jiao, F.; Xu, B. The

Central Role of Bicarbonate in the Electrochemical Reduction of Carbon Dioxide on Gold. J. Am. Chem. Soc. 2017, 139, 3774-3783.

25



24.

Page 26 of 28

Chen, Y.; Li, C. W.; Kanan, M. W. Aqueous CO2 Reduction at Very Low Overpotential

on Oxide-Derived Au Nanoparticles. J. Am. Chem. Soc. 2012, 134, 19969-19972. 25.

Wuttig, A.; Liu, C.; Peng, Q.; Yaguchi, M.; Hendon, C. H.; Motobayashi, K.; Ye, S.;

Osawa, M.; Surendranath, Y. Tracking a Common Surface-Bound Intermediate During CO2-toFuels Catalysis. ACS Cent. Sci. 2016, 2, 522-528. 26.

Shrikrishnan, S.; Sankaran, K.; Lakshminarayanan, V. Electrochemical Impedance

Analysis of Adsorption and Enzyme Kinetics of Calf Intestine Alkaline Phosphatase on SamModified Gold Electrode. J. Phys. Chem. C. 2012, 116, 16030-16037. 27.

Lee, S.; Hong, S.; Lee, J. Bulk pH Contribution to CO/HCOO− Production from CO2 on

Oxygen-Evacuated Cu2O Electrocatalyst. Catal. Today. 2016, 288 (supplement C), 11-17. 28.

Innocent, B.; Liaigre, D.; Pasquier, D.; Ropital, F.; Léger, J.-M.; Kokoh, K. B. Electro-

Reduction of Carbon Dioxide to Formate on Lead Electrode in Aqueous Medium. J. Appl. Electrochem. 2008, 39, 9658-9664. 29.

Kortlever, R.; Tan, K. H.; Kwon, Y.; Koper, M. T. M. Electrochemical Carbon Dioxide

and Bicarbonate Reduction on Copper in Weakly Alkaline Media. J. Solid State Electrochem. 2013, 17, 1843-1849. 30.

Hori, Y.; Suzuki, S. Electrolytic Reduction of Bicarbonate Ion at a Mercury Electrode. J.

Electrochem. Soc. 1983, 130, 2387-2390. 31.

Narayanan, S. R.; Haines, B.; Soler, J.; Valdez, T. I. Electrochemical Conversion of Carbon

Dioxide to Formate in Alkaline Polymer Electrolyte Membrane Cells. J. Electrochem. Soc. 2011, 158, A167-A173. 32.

Feng, X.; Jiang, K.; Fan, S.; Kanan, M. W. A Direct Grain-Boundary-Activity Correlation

for CO Electroreduction on Cu Nanoparticles. ACS Cent. Sci. 2016, 2, 169-174.

26


Page 27 of 28 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


33.

Li, C. W.; Ciston, J.; Kanan, M. W. Electroreduction of Carbon Monoxide to Liquid Fuel

on Oxide-Derived Nanocrystalline Copper. Nature 2014, 508, 504-507. 34.

Schouten, K. J. P.; Kwon, Y.; van der Ham, C. J. M.; Qin, Z.; Koper, M. T. M. A New

Mechanism for the Selectivity to C1 and C2 Species in the Electrochemical Reduction of Carbon Dioxide on Copper Electrodes. Chem. Sci. 2011, 2, 1902-1909. 35.

Sheng, H.; Oh, M. H.; Osowiecki, W. T.; Kim, W.; Alivisatos, A. P.; Frei, H. Carbon

Dioxide Dimer Radical Anion as Surface Intermediate of Photoinduced CO2 Reduction at Aqueous Cu and Cdse Nanoparticle Catalysts by Rapid-Scan Ft-Ir Spectroscopy. J. Am. Chem. Soc. 2018, 140, 4363-4371. 36.

Calle-Vallejo, F.; Koper, M. T. M. Theoretical Considerations on the Electroreduction of

CO to C2 Species on Cu(100) Electrodes. Angew. Chem., Int. Ed. 2013, 52, 7282-7285. 37.

Montoya, J. H.; Peterson, A. A.; Nørskov, J. K. Insights into C-C Coupling in CO2

Electroreduction on Copper Electrodes. ChemCatChem. 2013, 5, 737-742. 38.

Kas, R.; Kortlever, R.; Yılmaz, H.; Koper, M. T. M.; Mul, G. Manipulating the

Hydrocarbon Selectivity of Copper Nanoparticles in CO2 Electroreduction by Process Conditions. ChemElectroChem. 2015, 2, 354-358. 39.

Nie, X.; Esopi, M. R.; Janik, M. J.; Asthagiri, A. Selectivity of CO2 Reduction on Copper

Electrodes: The Role of the Kinetics of Elementary Steps. Angew. Chem., Int. Ed. 2013, 52, 24592462.

27



Page 28 of 28

TOC Graphic.

28


Revisiting Electrochemical Reduction of CO2 on Cu Electrode: Where

Recommend Documents