Chong
Wang
a,
Chengdeng
Wang
a,
Zhihao
Xiong
a,
Jiashuai
Wang
a,
Wenyuan
Zhang
a,
Haofeng
Shi
a,
Donghua
Wang
a,
Yousong
Gu
a,
Zhiming
Bai
*c,
Yan
Gao
*b and
Xiaoqin
Yan
*a
aSchool of Materials Science and Engineering, University of Science and Technology Beijing, Beijing 100083, People's Republic of China. E-mail: xqyan@mater.ustb.edu.cn
bLaboratory of Nanomaterials, National Center for Nanoscience and Technology, Beijing, 100190, People's Republic of China. E-mail: gaoyan@nanoctr.cn
cSchool of Civil and Resource Engineering, University of Science and Technology Beijing, Beijing 100083, P. R. China. E-mail: baizhiming2008@126.com
First published on 20th May 2022
Electrochemical CO2 reduction reactions (CO2RRs) have been recognized as a promising solution to environmental and energy problems due to their ability to convert intermittent renewable electricity into hydrocarbon fuels and value-added chemicals. Copper has a unique ability to electrochemically reduce CO2 to produce C2+ hydrocarbon products compared to other metals. However, it is still suffering from low Faraday efficiency (FE) limitations for specific products. Herein, we have prepared an Ag-modified Cu/CuO–Ag catalyst with a porous structure via a galvanic replacement reaction method, which reduced CO2 to C2+ products with the FE of 52.5% at −1.1 V (vs. RHE). The experimental results and density functional theory (DFT) calculations show that the addition of Ag in Cu is beneficial to increasing the coverage of *CO on the Cu surface, which can decrease the potential barrier energy of the C–C coupling reaction and favor the generation of C2+ products. The results of this study may be helpful to the design of efficient tandem catalysts for electrochemical CO2 reduction.
Nowadays, the common catalysts for the CO2RR are based on transition metals or their compounds, which have been reported to be used for electrochemical CO2 reduction to produce different products.10,11 For example, Au, Ag and Zn are capable of reducing CO2 to carbon monoxide, while Sn, In and Bi are more inclined to reduce CO2 to formate.12–18 Copper has attracted great interest as the only metal material capable of electrochemically converting carbon dioxide to hydrocarbons and alcohols in aqueous solution.19–21 Unfortunately, Cu-based catalysts still face a series of issues in the reduction of CO2, such as high overpotential, low current density and low product selectivity.22 Several strategies have been proposed to improve the properties of Cu-based catalysts, including engineering low-coordinated sites and defects,23–25 modifying the size and shape of nanostructures,26,27 and tuning the oxidation state.22,28,29 For example, it has been reported that a copper catalyst synthesized from the reduction of oxidized Cu showed higher activity and better C2+ product selectivity in the electrochemical reduction of carbon dioxide compared with unoxidized copper foil.30 The results show that the reduced copper oxide electrode can expose more optimized surfaces with edges and steps for the initial activation of carbon dioxide molecules and the dimerization of related intermediates.
Currently, a promising method is to combine copper-based catalysts with second catalysts (such as gold, silver and zinc) to further increase the surface coverage of *CO, which contributes to the formation of C2+ products.31–33 For example, Du et al. constructed silver–copper bimetallic catalysts with sharp interfaces via sequential precipitation for the efficient production of ethylene. The Faraday efficiency of ethylene production was more than twice that of the pure copper catalyst. DFT calculation shows that the adsorption energy of CO on Cu is three times that of Ag. The CO molecules are first formed on Ag atoms and then trapped by Cu atoms through the Ag/Cu interface, resulting in a large number of CO molecules aggregated on the Cu side of the Ag/Cu interface. Therefore, the increase of the FE for the C2+ products is due to the tandem effect of the silver–copper metal sites near the interface.34 Although CuAg bimetallic catalysts can improve the C2+ product selectivity, the activity is still low and the current density is not sufficient for commercialization.35 Moreover, the reaction mechanism of converting CO2 to C2+ products over Ag-modified Cu-based catalysts should be further elucidated.
Here, Ag-modified Cu/CuO (Cu/CuO–Ag) catalysts with a porous 3D structure were synthesized by a galvanic replacement reaction (GRR) between Ag+ and Cu. This special mesoscopic architecture provides the best combination between copper and silver, which dramatically improves the Faraday efficiency and current density of ethylene formation. The electrochemical studies show that Cu/CuO–Ag has good catalytic activity toward the reduction of CO2, and the FE for C2+ reaches 52.5% at −1.1 V (vs. RHE), much greater than that of the Cu/CuO catalyst (32.2%). DFT calculation shows that increasing the coverage of *CO on the Cu surface facilitates the C–C dimerization reaction and improves the selectivity of C2+ products.
(1) |
To evaluate the selectivity of CO2 reduction, we used online electrochemical gas chromatography (GC, Shimadzu GC-2014c) to detect the gas product during the CO2RR through the continuous flow of an electrolytic cell. The GC was calibrated regularly using a standard gas combination under standard circumstances. The faradaic efficiency (FE) data of the gas product was computed from the following equation:
(2) |
The electrolytes were collected separately after 1 h of electrolysis at each potential and the liquid products were analyzed quantitatively using the Ascend-400 MHz nuclear magnetic resonance (NMR) spectrometer. After 1 h of electrolysis, 0.5 mL of electrolyte was removed, and 0.1 mL of D2O and 1.67 ppm of dimethyl sulfoxide (DMSO) were added to it as an internal standard. A pre-saturation technique is used to suppress the water signal. The FE data for liquid products are computed using the following formula:
(3) |
We first optimized the crystal structure and found the equilibrium lattice constant to be αCu = 3.631. Then a 4-layer model with a p (3 × 3) Cu(111) supercell was used, with the top two layers relaxed and the bottom two layers fixed. The Nudged Elastic Band (NEB)41 and Dimer's methods42 were used to locate the transition state structures, and to further optimize the transition state using a residual force convergence criterion of less than 0.05 eV Å−1.
HER (proton reduction): 2H+ + 2e− → H2 | (4) |
HER (water splitting): 2H2O + 2e− → H2 + 2OH− | (5) |
Metal deposition: Cu2+ + 2e− → Cumet | (6) |
The morphological characteristics of copper foams after deposition for 20 s are shown in Fig. S1 (ESI†). The original Cu foam shows an obvious pore size gradient along the surface normal. The smallest pore is close to the supporting electrode (copper foil), and the largest pore is located on the outermost surface of the Cu foam. The interconnected pores are visible with surface pore sizes varying from 50 to 80 μm, which provide adequate space for the electrolyte to circulate freely. The sidewalls of the pores are composed of loosely packed copper dendrites, thus forming “secondary” nanopores in the Cu foam. Then, the Cu foam was further thermally annealed in air at 250 °C for 3 h, thereby converting the catalyst materials into Cu/CuO foams. The overall morphology of the copper foam remains essentially unchanged following thermal annealing, which is proved by SEM (Fig. 2a and b). The Cu/Cu2O–Ag was prepared by immersing Cu/CuO into an aqueous AgNO3 solution for 1–5 min. Since the standard reduction potential of the Cu2+/Cu redox couple (0.34 V vs. SHE) is lower than that of Ag+/Ag (0.80 V vs. SHE), the following primary galvanic replacement reaction (GRR) occurs spontaneously:
(7) |
Fig. 2 (a) and (b) SEM images of Cu/CuO. (c) and (d) SEM images of Cu/CuO–Ag. (e–h) EDX images of overlap, Cu, O and Ag of Cu/CuO–Ag. |
After 3 min of GRR in 10 mM AgNO3 solution, silver nanoparticles (NPs) with the size of 50–100 nm were uniformly distributed on the surface of Cu/CuO foams (Fig. 2c and d). The EDX shows that Cu, O and Ag are uniformly distributed in the Cu/CuO–Ag nanocomposites (Fig. 1e–h), which matches well with the results of the above SEM characterizations. The TEM image (Fig. S2, ESI†) shows that silver nanoparticles are attached to copper dendrites. In addition, we also characterized the distribution of Ag nanoparticles in the Cu/CuO–Ag1 and Cu/CuO–Ag5 samples. The images in Fig. S3 (ESI†) show a gradual increase in the average size of the Ag nanoparticles from the Cu/CuO–Ag1 to Cu/CuO–Ag5 samples. The roughness factor of the foam structured catalyst was estimated using the double layer capacitance in the nonfaradaic regions of the CV method. The electrodeposited catalysts have a larger roughness factor than that of the Cu foil catalysts. The roughness of the catalysts follows the order as Cu foam < Cu/CuO–Ag < Cu/CuO (Fig. S4 and Table S1, ESI†).
The XRD of Cu/CuO and Cu/CuO–Ag show peaks that belong to the diffraction of CuO(PDF NO. 45-0793) and Cu(PDF No. 04-0836) (Fig. 3a). The peaks at 38.3°, 44.4°, 64.4° and 77.3° of Cu/CuO–Ag could be ascribed to Ag(111), Ag (200), Ag(220) and Ag(310), respectively. XPS was used to study the chemistry states and surface composition of the Cu/CuO–Ag catalysts. The catalyst exhibits a pair of spin–orbiting double Cu 2p peaks, which may be attributed to Cu 2p3/2 and Cu 2p1/2, respectively (Fig. 3b). The fitting peaks at 933.8eV and 953.2 eV are ascribed to the Cu 2p3/2 and Cu 2p1/2 peaks of Cu2+, respectively. Furthermore, the presence of CuO is supported by the presence of two discernible shake-up satellite peaks.43 The fitted peak at 932.1 eV makes it hard to distinguish the Cu+ and Cu0 directly by XPS since the difference in binding energy between Cu+ and Cu0 is only 0.1 eV.44 The chemical state of Cu was further verified using Auger electron spectroscopy (Fig. 3c). Cu LMM XAES spectra indicate the presence of two surface Cu species near 914 eV (Cu+) and 918 eV (Cu0). The silver 3d3/5 and 3d3/2 peaks of the Cu/CuO–Ag are located at 368.6 and 374.6 eV (Fig. 3d), which also confirms the existence of Ag on Cu/CuO–Ag.45 The Cu/CuO and Cu/CuO–Ag catalysts were also characterized after pre-reduction (Fig. S5, ESI†). It can be confirmed by XRD and XPS that the CuO components have been reduced to Cu in both catalysts.
Fig. 3 (a) XRD patterns of Cu/CuO and Cu/CuO–Ag. (b) High-resolution Cu 2p XPS spectra, (c) Auger spectra of Cu LMM, and (d) Ag 3d XPS spectra of Cu/CuO–Ag. |
The selectivity and activity of the CO2RR were evaluated in a three-electrode H-type cell at a constant external electrode potential (Fig. S6, ESI†). The catalytic activity of the prepared catalysts was first investigated by linear voltammetry. As shown in Fig. 4a, both catalysts exhibit higher current density to reduce CO2 compared to the copper foil. The total current density of Cu/CuO and Cu/CuO–Ag increases from 1.05 mA cm−2 to 18.42 mA cm−2 and from 1.45 mA cm−2 to 22.76 mA cm−2 when the potential decreased from −0.47 V to −1.20 V (vs. RHE). Therefore, the total current density of the Cu/CuO and Cu/CuO–Ag catalysts shows similar trends, but the current density of the Cu/CuO–Ag catalyst is greater than that of the Cu2O catalyst, which indicates that the introduction of Ag improved the catalytic activity of the catalysts. In addition, the total current density of Cu/CuO is greatly improved compared with the original Cu foam at different potentials (Fig. S8a, ESI†), which may be caused by the significant increase in surface roughness of the Cu foam after annealing reduction (Table S1, ESI†).
The selectivity of the products was obtained by GC and 1H NMR analysis after applying different potentials in a 0.1 M KHCO3 solution with saturated CO2 and electrolysis at each potential for 1 h. The geometric current density of Cu/CuO–Ag remained stable during 60 min of electrolysis (Fig. S7, ESI†). The FE values of the carbon dioxide reduction products on the Cu/CuO and Cu/CuO–Ag catalysts are shown in Fig. 4b and c. For the Cu/CuO–Ag catalysts, the reduction of CO2 produced almost no C2+ product, and the main product is H2 with FE of 62.01–56.7% at overpotentials from −0.6 V to −0.7 V (vs. RHE). Significantly, the FE of C2H4 increases from 13.9% at −0.8 V to 38.7% at −1.1 V (vs. RHE). The FE of C2H4 for the Cu/CuO–Ag catalyst was 38.7% at −1.2 V (vs. RHE), which is greater than the Faraday efficiency of Cu/CuO (22.4%) at the identical potential. By contrast, copper foil and copper foam catalysts produce much fewer C2H4 (Fig. S8b, ESI†). The copper foam catalyst generates few hydrocarbons and the major formic acid product exceeds 26% at −1.0 V (vs. RHE), which is consistent with the literature.46 Quite intriguingly, the Faraday efficiency of the formate for the Cu/CuO catalyst has remained at a relatively low level and would not exceed 11% at any electrolytic potential. In contrast, the Faraday efficiencies for ethylene and ethanol products reach 22.4% and 12.8%, respectively.
The key to reducing CO2 to HCOOH or CO is whether the CO2 is adsorbed by its oxygen atoms or carbon atoms. There is no breaking of the CO bond during the formation of HCOOH. The CO2 molecules form *CO2˙− free radical anions on the surface of the catalysts and extract a proton from the adjacent water, which then results in the production of formate.47 The different distribution of CO2RR products between Cu foams and Cu/CuO catalysts is mainly due to the weakly coordinated surface sites caused by electrochemical reduction of Cu/CuO foams under CO2RR conditions, which are particularly active for carbon–carbon coupling reactions.48
Fig. 4d summarizes the C2+ product FEs for Cu/CuO and Cu/CuO–Ag catalysts. The Ag/Cu catalysts show the highest FEs for the C2+ products (52.5% at −1.1 V vs. RHE) as compared to Cu/CuO (32.2%), which indicates that the addition of Ag nanoparticles can efficiently improve the selectivity of the C2+ products. We also test the electrochemical reduction characteristics of the Cu/CuO–Ag catalysts by varying the Ag content. FE of the C2+ product shows a volcanic variation with increasing Ag loading, which indicates that there is an optimal amount of Ag loading (Fig. S9, ESI†). When the content of Ag is low, it does not provide enough CO intermediates for the copper atoms to promote the C–C coupling reaction. While the higher content of Ag leads to severe aggregation of silver nanoparticles, resulting in elevated FE of CO and lower FE of C2+. Ag foam and Cu–Ag foam were also prepared in a similar way (Fig. S10, ESI†). XRD patterns show that separated pure metal Cu and Ag phases are also formed in the Cu–Ag foam (Fig. S11, ESI†). Their performance for CO2 reduction was tested under the same conditions (Fig. S12, ESI†), and the CO2 reduction product on the Ag foam is dominated by CO, yielding a very small amount of C2H4 (FE < 3%). In contrast, the selectivity of the Cu–Ag foam for the C2+ product is significantly improved, reaching a Faraday efficiency of 27.5% for the C2+ product at −1.1 V (vs. RHE). However, it is lower than that of Cu/CuO–Ag (52.5%), indicating that the synergistic effect between the oxidation-derived Cu and Ag modifications leads to the high selectivity of the Cu/CuO–Ag catalyst for the C2+ product.
Cu/CuO–Ag shows good stability with continuous reduction of CO2 to C2H4 at −1.1 V (vs. RHE) for 8 h (Fig. 4e). The FE 39% of the C2H4 products remains almost unchanged. A longer test was conducted for 20 h, and the catalytic activity of Cu/CuO–Ag and the faradaic efficiency for ethylene hardly changes (Fig. S13, ESI†). In addition, the Cu/CuO–Ag catalyst remains in its original porous morphology (Fig. S14, ESI†), which indicates that the Cu/CuO–Ag catalyst has excellent stability.
The FE of C2H4 for the Cu/CuO–Ag catalyst is obviously higher than that of the Cu/CuO catalyst, which indicates that the introduction of Ag contributes to the production of C2+ products. Furthermore, we discuss the possible catalytic mechanism of reducing CO2 to C2H4 in this catalyst (Fig. 5a). At present, *CO is recognized as a critical intermediary in the electrochemical reduction of carbon dioxide to hydrocarbons.49 The formation of *CO intermediates from CO2 molecules usually undergoes two processes. First, the CO2 molecule is caught by Cu and Ag atoms, and both a proton and an electron are transferred.50
(11) |
Next, the *COOH intermediate is further reduced to *CO.
(12) |
Ag atoms are more active than Cu atoms in the production of carbon monoxide molecules.51 The *CO generated on Ag is more easily desorbed as *CO is weakly bound to Ag, increasing the local CO concentration around the electrode surface. The CO will be temporarily trapped in the pores of the foam catalyst and re-adsorbed by the Cu sites. Additionally, the *CO intermediate produced on the Ag atoms can be captured by the Cu atoms through the Cu–Ag interface, leading to a high surface coverage of *CO. In the step of further reduction of the CO intermediate to ethylene on the Cu surface, the rate-determining step (RDS) is the dimerization of CO leading to the formation of a C–C bond.52 We consider that the coverage of CO intermediates on the catalyst surface has an important effect on the CO dimerization reaction and high surface coverage of *CO may facilitate the synthesis of the C2+ product, which is further proved by DFT calculation below.
Computationally, a (3 × 3 × 4) Cu(111) supercell was built to simulate the catalyst. With DFT, we calculated the C–C coupling barrier energy for *CO coverage of 2/9 ML and 3/9 ML to simulate the effect of high and low* CO coverage on the Cu surface for the synthesis of C2+ products (Fig. S15, ESI†), respectively. Although solvent molecules have a certain effect on the CO2RR,53 in order to simplify the model, we ignore the role of solvent molecules, and here the geometry of *OCCO is stable, the same as that under no charged water layer.52 As shown in Fig. 5b and c, increasing the CO coverage from 2/9 ML to 3/9 ML decreases the barrier from 1.62 eV to 1.13 eV, a reduction of 0.49 eV. The lowering of the barrier energy of the C–C coupling reaction is due to the interaction between the adsorbates (CO), which weakens the carbon bonds on the surface and makes the carbon–carbon bonds more easily formed (Table S2, ESI†).
Footnote |
† Electronic supplementary information (ESI) available. See DOI: https://doi.org/10.1039/d2ma00188h |
This journal is © The Royal Society of Chemistry 2022 |