Titania-Modified Silver Electrocatalyst for Selective CO2 Reduction to

Subscriber access provided by UNIV OF DURHAM

Article

Titania-Modified Ag Electrocatalyst for Selective CO Reduction to CHOH and CH from DFT Study 2

3

4

Lina Zhai, Chaonan Cui, Yuntao Zhao, Xinli Zhu, Jinyu Han, Hua Wang, and Qingfeng Ge J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.7b03314 • Publication Date (Web): 06 Jul 2017 Downloaded from http://pubs.acs.org on July 13, 2017

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.

The Journal of Physical Chemistry C 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

The Journal of Physical Chemistry

Titania-Modified Ag Electrocatalyst for Selective CO2 Reduction to CH3OH and CH4 from DFT Study Lina Zhai, † Chaonan Cui, † Yuntao Zhao, † Xinli Zhu, † Jinyu Han, † Hua Wang*, † and Qingfeng Ge*, †, ‡ †

Collaborative Innovation Center of Chemical Science and Engineering, Key Laboratory for Green Chemical Technology, School of Chemical Engineering and Technology, Tianjin University, Tianjin 300350, China

‡

Department of Chemistry and Biochemistry, Southern Illinois University, Carbondale, Illinois 62901, United States

Corresponding Authors *E-mail: [email protected]. Tel.: 86-139-2069-3419 (H.W.) *E-mail: [email protected]. Tel.:1-618-453-6406 (Q.G.).

ACS Paragon Plus Environment

1


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

ABSTRACT

Electrochemical reduction of CO2 to produce useful fuels and chemicals is one of the attractive means to reuse CO2. Herein, we constructed a (TiO2)3/Ag(110) model electrocatalyst and examined CO2 reduction pathways. Our results show that the interface between oxide and supporting Ag provides the active sites for CO2 adsorption and activation. These active sites enable the electron transfer to the adsorbed CO2. In this setup, Ag acts as an electron donor, partially reducing the supported (TiO2)3 and supplies the needed electrons to the adsorbed CO2. Once CO2* is formed at the interface, the subsequent hydrogenation steps take place sequentially. Our results further indicate that the dominating pathway to produce CH3OH is via the H2COOH* intermediate following the formation of HCOO*. The formation of H2COOH* with a free energy of 0.47 eV is the potential-limiting step. Furthermore, protonating H2COOH* followed by dehydration to CH3O* and hydrogenation of CH3O* leads to CH4 formation. The COOH* pathway may converge to the same H2COOH* intermediate instead of forming CO*. These results demonstrated the benefit of metal supported metal oxides as electrocatalysts to produce CH3OH or CH4 from electrochemical reduction of CO2.


2

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


1. INTRODUCTION Climate change caused by greenhouse gas emission and accumulation has become one of the greatest threats to a sustainable development of human society.1 Among the greenhouse gases, carbon dioxide (CO2) from anthropogenic activities contributes most to the increase of the atmospheric and oceanic CO2 levels. Over the past decades, there have been increasing efforts to address the CO2 issue, including carbon capture and sequestration as well as converting captured CO2 to fuels and useful chemical products.2 Among various CO2 conversion approaches,3−6 electrochemical reduction of CO2 has the advantage of moderate reaction conditions and diverse products, depending on applied potential and catalysts.7,8 In addition, electrochemical CO2 reduction provides an alternative to store energy from the intermittent renewable sources like solar and wind, making those sources more sustainable.9 However, electrochemical CO2 reduction faces many challenges, including high overpotential, poor product selectivity and low energy efficiency. Therefore, a better understanding of the CO2 reduction mechanism will help to design highly efficient catalysts for electrochemical reduction of CO2 at a practical level. The electrochemical reduction of CO2 involves a complicated multiple proton−electron transfer process. Several multi-electron reduction products have been identified on different catalysts.9,10 The thermodynamic redox potentials with reference to the reversible hydrogen electrode (RHE) for selected C1 products11 are listed below: 2H+ + 2e− → H2

E° = 0.00 V

CO2 + e− → CO2−

E° = −1.90 V

CO2 + 2H+ + 2e− → CO + H2O

E° = −0.11 V

CO2 + 2H+ + 2e− → HCOOH

E° = −0.20 V

CO2 + 6H+ + 6e− → CH3OH + H2O

E° = +0.03 V


3


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

CO2 + 8H+ + 8e− → CH4 + 2H2O

Page 4 of 27

E° = +0.17 V

Thermodynamically, the reduction to both CH3OH and CH4 are favorable with standard potentials of +0.03 V (RHE) and +0.17 V (RHE), respectively. However, these are generally minor products on most electrodes, and even if they form, the overpotentials are high.9 On many electrodes, it has been believed that one-electron reduction of CO2 to CO2−, with a potential of −1.90 V (RHE), is the potential limiting step.12−14 Subsequent product formation depends on how efficiently CO2− can be formed and further reduced under the reaction conditions. Earlier studies on CO2 electrochemical reduction have been focused on metallic electrodes.13,15 Ag is one of the widely-used electrode materials and has been known to produce mainly CO.13 At a potential less than −1.2 V (RHE), multi-electron reduction products such as CH3OH and CH4 have been reported on Ag surface but their quantities are minor compared to CO.9,16 Both structural and compositional modifications have been proposed to improve the activity and selectivity.7,8,17−20 Novel materials, including 2D materials, metal organic materials, and metal oxides have been explored for CO2 electrochemical reduction.21−23 Ma et al. used silver supported on TiO2 for CO2 electroreduction and reported a 73 mV lower onset potential for CO2− formation than that on the Ag nanoparticles.24 They suggested that the presence of TiO2 facilitated CO2 adsorption and stabilized the CO2− intermediate. In fact, TiO2 may serve as a redox electron carrier to reduce CO2 while Ag assists in the formation of the final product.24 Particularly, He and co-workers reported that charge transfer can be enhanced when CO2 is adsorbed on metal oxides.25 In addition to metal-based electrocatalysts, metal oxides, including TiO2, Cu2O, RuO2, as catalysts for CO2 electrochemical reduction also showed impressive performances for CH3OH and CH4 production.26−33 Ramesha et al. suggested that partially reduced metal oxides, i.e. the


4

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


Ti3+ sites formed on the nanostructured TiO2 surface are catalytically active and responsible for charge transfer to CO2 and subsequent reduction reactions.28 Furthermore, multi-electron reduction products such as CH3OH and CH4 are more likely produced due to the preferential formation of HCOO* on the rutile surfaces.31,33 However, electron conductivity and the stability of the oxides as cathode materials are challenges for the metal oxide-based electrodes. As an alternative to using either metals or metal oxides as electrode materials, metal/metal oxide composites could exhibit higher activity for CO2 reduction with diverse products with good electrode conductivity.34−37 However, the active oxides may have to be regenerated periodically to replenish the losses during the reaction. In the present study, we constructed a model oxide−metal electrocatalyst by supporting a (TiO2)3 cluster on the Ag(110) surface and examined the reaction pathways and possible products based on density functional theory calculated energetics. Although the actual catalytic performance may depend on the size distribution of the supported TiO2 particles, this simplified model is expected to provide an understanding of the role of those highly undercoordinated rim sites at the oxide/metal interface in CO2 activation and reduction. In fact, our results showed the benefit of the metal oxide−metal composites as catalysts for CO2 electrochemical reduction: the presence of the active metal oxide enables the electron transfer and reduces the potential for the reduction of CO2. This result provides a novel approach to design more efficient catalysts for electrochemical reduction of CO2. 2. COMPUTATIONAL DETAILS All calculations were performed using the Vienna ab initio simulation package.38,39 The effective core potential is described by the projector augmented wave method,40 and the exchange correlation energy is calculated using the Perdew−Burke−Ernzerhof (PBE) functional.


5


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

A plane wave basis set with a cutoff energy of 400 eV was used to describe the valence electrons.41 Similar set of parameters was used in previous studies and found to provide converged structures and energies.42,43 A slab model of Ag(110) was constructed based on the optimized lattice constant of 4.15 Å. The slab consists of four atomic layers of Ag atoms and is separated by a vacuum gap of 12 Å. A (3 × 4) surface unit cell resulted in a supercell of 12.46 Å × 11.75 Å × 16.41 Å. A k-points grid of 4 × 4 × 1 was used to sample the Brillouin-zone. In our calculations, the atoms in the bottom two layers were fixed at their bulk position and those in the top two layers together with the supported TiO2 cluster and the adsorbates were allowed to relax. We constructed the model system by supporting a (TiO2)3 cluster on the Ag(110) surface, as shown in Figure 1. Our results indicated that the most stable (TiO2)3 cluster (Figure 1a) has a Cs symmetry. Possible binding structures were tested and the most stable one was shown in Figure 1b. As shown in Figure 1b, the cluster is anchored to the surface through three O atoms (O1, O2, O3) with no direct bonding between other atoms of the cluster and the surface atoms.


6

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


Figure 1. Structure of (TiO2)3 cluster (a) and side and top views of (TiO2)3/Ag(110) (b). Three regions on (TiO2)3/Ag(110) for possible CO2 adsorption are also divided in the side view. Ag atoms are in blue, Ti atoms in green and O atoms in red. Different O atoms are labeled with numbers.

The binding energy of the cluster on the Ag surface, defined as the energy difference between the slab with (TiO2)3 and the sum of the isolated cluster and the Ag(110) slab, is −1.41 eV. Bader charge analysis indicates that a net charge of −0.65 |e| was transferred to the cluster from the Ag slab. Bader charge analysis (Table S1 of Supporting Information) showed that the net charge on the Ti1 atom was lower than the charge of Ti atom in bulk TiO2, indicating that Ti was partially reduced in the supported (TiO2)3 cluster. The following equation was used to calculate the adsorption energy of an adsorbate ∆Eads = E(adsorbate/(TiO2)3/Ag(110)) – E(adsorbate) – E((TiO2)3/Ag(110))

(1)

where E(adsorbate/(TiO2)3/Ag(110)) is the total energy of an adsorbate bound to (TiO2)3/Ag(110), E(adsorbate) is the total energy of the isolated adsorbate, and E(TiO2)3/Ag(110)


7


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

is the total energy of bare (TiO2)3/Ag(110).44 A negative value of ∆Eads corresponds to an exothermic adsorption process. The free energy diagrams for CO2 reduction along different pathways were calculated with reference to the computational hydrogen electrode (CHE) proposed by Nørskov et al.45 According to this method, the proton-coupled electron transfer step has a negligible kinetic barrier at room temperature and the step with the most positive energy difference would be the rate-determining, or the potential-limiting step.11 Consequently, the free energy of each species can be obtained from the following equation G = Eelec + ZPE – TS

(2)

where G is the free energy, Eelec is the electronic energy of the species from the DFT calculation, ZPE is the zero-point energy, T is the temperature and S the entropy. ZPE and vibrational contribution to the entropy of the adsorbed species can be obtained using the calculated vibrational frequencies.46 The chemical potential of a proton−electron pair (µ[H+ + e−]) at U = 0 V equals to half of hydrogen (0.5 µ[H2]) at 101 325 Pa of H2, 298 K, and all pH values. Thus, free energy change of each elementary step at external potential U will be shifted by |e|U as ∆G (U) = ∆G (U = 0) + |e|U

(3)

where |e| is the electronic charge. According to Nørskov et al., the thermodynamic activation barrier equals the largest of the free energy differences.45 3. RESULTS AND DISCUSSION 3.1 CO2 adsorption and activation


8

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


The adsorption and activation of CO2 is considered as the first step for CO2 reduction. In this section, stable CO2 adsorption configurations on (TiO2)3/Ag(110) were determined and shown in Figure 2. Adsorption energies and results of Bader charge analysis are listed in Table 1.

Figure 2. Stable CO2 adsorption structures on (TiO2)3/Ag (110). (a) The C atom binds Ag(110) and OI atom binds Ti of (TiO2)3; (b) The C atom binds Ag(110) and both OI and OII atoms bind Ti of (TiO2)3; (c) The OI atom binds Ti of (TiO2)3 and OII points to Ag(110) with C bending away from the surface. Bond lengths are in Å. Ag atoms are in blue, Ti atoms in green, O atoms in red and C atoms in gray. O atoms of CO2 in different positions are labeled with Roman numerals I and

Figure 2a−c show the optimized CO2 adsorption structures with CO2 bridging (TiO2)3 and Ag(110). As shown in Figure 2a, CO2 interacts with Ag through its C atom and bonds with the unsaturated Ti atom through the OI atom. The adsorbed CO2 is stabilized in a bent conformation with an O−C−O angle of 123.8° and elongated C−O bonds (1.21 Å, 1.35 Å), as compared with that in the isolated CO2 molecule (1.18 Å). Density of state (DOS) analysis (Figure S1, S2 in Supporting Information (SI)) confirms that C and Ag form a new bond, corresponding to the overlap of C 2p states and Ag 4d states. Figure 2b shows the adsorption structure where both O atoms of CO2 bind to the same Ti atom. CO2 can also bridge the cluster and surface through both of its oxygen atoms, forming a structure shown in Figure 2c. However, an OII−Ag distance of 2.61 Å indicates that the interaction is not likely strong. Comparing ∆Eads shown in Table 1, structures a and b are significantly more stable than c. We selected the most stable (a) as the starting structures for further reduction. In addition, CO2 adsorption on Ag(110) away from


9


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

(TiO2)3 is rather weak, similar to that on the Ag(110) surface, with ∆Eads = −0.02 eV. In fact, adsorbed CO2 on Ag(110) remains in a nearly linear configuration. The structures of CO2 on the (TiO2)3 without direct interaction with Ag(110) are shown in Figure S3 in Supporting Information. In Figure S3a, CO2 is adsorbed in a bidentate configuration with O atom binding Ti and the C atom with the O atom of (TiO2)3, similar to adsorbed CO2 on the metal oxide surfaces.47 Figure S3b shows that the adsorbed CO2 is activated into a carbonatelike structure, binding the Ti atom in a chelate structure. Although their adsorption energies (−0.59 eV and −1.84 eV, respectively) indicate that both structures are favorable thermodynamically, their reduction pathways are expected to be the same as those on the surface of TiO2.25,

48

Therefore, we did not examine those pathways and only focus on the steps

occurring at the TiO2/Ag interface. Bader charge analyses were conducted for the adsorption structures and the results are listed in Table 1. There is a net charge of −0.80 |e| and −1.00 |e| transferred to the adsorbed CO2 in structures a and b, respectively. Previous studies showed that the first electron transfer to CO2 with an equilibrium potential of −1.90 V (RHE) is the potential-limiting step for CO2 electroreduction.13 Clearly, the presence of TiO2 on Ag(110) creates active sites for CO2 adsorption. Furthermore, the adsorption of CO2 also oxidized the Ti1 atom and increased its charge by 0.22 |e| from the Ti atom in the bare (TiO2)3/Ag(110) (see Table S1 in Supporting Information). The results indicate that in addition to providing the active site for CO2 adsorption and activation, the (TiO2)3 cluster also acts as an intermediary to facilitate electron transfer from Ag to adsorbed CO2. Notably, when the reduction takes place on (TiO2)3/Ag(110), both (TiO2)3 and Ag(110) contribute to the charge transferred to CO2 although a major portion of that is from


10

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


Ag(110). This is of particular importance for CO2 electrochemical reduction since catalysts such as supported TiO2 can promote electron transfer and improve current efficiency.

Table 1. DFT-PBE adsorption energies and total Bader charges of adsorbed CO2, (TiO2)3, and Ag(110) slab for CO2 on (TiO2)3/Ag(110) in configurations shown in Fig. 2

Structures

∆Eads (eV)

(TiO2)3/Ag(110)

Net Charge (|e|) Cluster

CO2

Ag(110)

−−

−0.65

−−

0.65

Fig. 2(a)

−1.10

−0.44

−0.80

1.24

Fig. 2(b)

−1.04

−0.41

−1.00

1.41

Fig. 2(c)

−0.63

−0.46

−0.54

0.99

3.2 CO2 reduction pathways We use the activated CO2 adsorption configuration shown in Figure 2a (denoted as CO2*) as the starting point for further reduction. The first hydrogenation of CO2* may occur at either the O or C atom, forming carboxyl (COOH*) or formate (HCOO*) species, respectively. Herein, we investigated the routes through both COOH* and HCOO* intermediates. The free energy diagrams for CO2 reduction on (TiO2)3/Ag(110) along the COOH* and HCOO* pathway at 0 V (RHE) are shown in Figure 3.


11


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

Figure 3. Free energy diagrams for CO2 reduction on (TiO2)3/Ag(110) along the COOH* (black lines) and HCOO* pathway (red lines) at 0 V (RHE), including the structures of COOH*, HCOOH*, HCOO*, H2COO* and H2COOH*. Ag atoms are in blue, Ti atoms in green, O atoms in red, H atoms in white and C atoms in gray.

3.2.1 The COOH* pathway On (TiO2)3/Ag(110), COOH* is formed through protonating the O atom of adsorbed CO2. This is different from CO2 reduction on pure Ag(110), shown in Figure S4 for comparison. On Ag(110), CO2 reduction follows CO2 → COOH* → CO* path with the formation of COOH* as the most endothermic step.49 Since CO2 binds Ag(110) weakly, the formation of COOH* follows a mechanism similar to the Eley−Rideal mechanism in gas−solid catalysis: CO2 reacts with the reduced proton on the Ag surface. In contrast, CO2 binds strongly at the TiO2/Ag(110) interface, making COOH* formation from CO2* much less endergonic than on Ag(110). The change of COOH* formation mechanism also switched the relative position of COOH* and CO*, making CO* formation from COOH* endergonic on (TiO2)3/Ag(110), with a free energy change of


12

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


0.69 eV. Clearly, the formation CO2* at the (TiO2)3/Ag(110) interface facilitates the electron transfer to adsorbed CO2 and promotes electrochemical CO2 reduction. Furthermore, CO* from COOH* is stable on (TiO2)3/Ag(110) as its desorption is endergonic with a free energy change of 0.46 eV. Consequently, CO* may be further reduced on (TiO2)3/Ag(110) to other products. Therefore, we also mapped out the reduction pathway from CO* and presented the free energy profile of the corresponding intermediates, including CHO*, CH2O* and CH3O*, in Figure 3. As shown in Figure 3, the process after CO* formation is downhill in free energy, in contrast to that on a metal surface such as Cu.50 The structures of those intermediates on (TiO2)3/Ag(110) are shown in Figure 4. All intermediates, including CO*, CHO* and CH2O*, adsorb through both C and O atoms at the metal−oxide interface. The distances between C atom and the Ag atom gradually increase with the increasing degree of hydrogenation of C atom. As shown in Figure 4d, the Ag−C in CH2O* is broken upon formation of CH3O*. On the other hand, the high oxygen affinity of (TiO2)3/Ag(110) keeps the O atom of CH3O* bound to Ti strongly. And this, in turn, will cause the C−O bond of CH3O* to break and eventually produce CH4 in the subsequent hydrogenation step.


13


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

Figure 4. Structures of reduced intermediates following the COOH* pathway of CO2 reduction on (TiO2)3/Ag(110). The intermediates are (a) CO*, (b) CHO*, (c) CH2O* and (d) CH3O*, respectively. Bond lengths are in Å. Ag atoms are in blue, Ti atoms in green, O atoms in red, H atoms in white and C atoms in gray.

These results also indicate that hydrogenation of COOH* on (TiO2)3/Ag(110) occurs more likely at the C-end of the intermediate, resulting in HCOOH* rather than CO* and H2O. This step has only a free energy barrier of 0.09 eV. The adsorbed formic acid species can be further hydrogenated to H2COOH*, which will then converge to the HCOO* pathway to produce CH3OH or CH4, as we will discuss in the next section. 3.2.2 The HCOO* pathway Another pathway of hydrogenating CO2* leads to the HCOO* intermediate. As shown in Figure 3, the HCOO* formation is thermodynamically preferable over COOH* on (TiO2)3/Ag(110). Following HCOO* formation, subsequent hydrogenation to H2COO* to H2COOH* are endergonic with free energy barriers of 0.32 eV and 0.47 eV, respectively. Further hydrogenation to CH3OH and CH4 are exergonic. Thermodynamically, the formation of CH3OH and CH4


14

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


through H2COOH* are both downhill in free energy. The formation of CH3OH is more favorable than CH4, by 0.10 eV. The most endergonic step along this pathway is the formation of H2COOH*, which will be the potential-limiting step for CO2 reduction on (TiO2)3/Ag(110). Comparing with the COOH* pathway, which has a limiting potential of 0.69 eV, the HCOO* pathway provides a lower energy channel to produce CH3OH or CH4. Based on the free energy consideration, the HCOO* pathway via the intermediate H2COOH* is preferable on (TiO2)3/Ag(110). We also depict the free energy diagram with an external potential of −0.47 V (RHE), as shown in Figure 5. The potential is selected based on the potential-limiting step, i.e. the formation of H2COOH*. Under this applied potential (U = −0.47 V vs RHE), the free energy change of all steps involving the proton−electron pairs will be corrected according to eq. 3. Under this potential, all steps following H2COOH* formation become downhill on the free energy diagram.

Figure 5. Free energy diagram for CO2 reduction on (TiO2)3/Ag(110) along HCOO* pathway at −0.47 V (RHE).


15


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

Along the HCOO* pathway, H2COOH* is the key intermediate and plays a central role in the entire mechanism. The role of H2COOH* can be illustrated in the cyclic process shown in Figure 6. Further hydrogenating H2COOH* is likely to break one of the two C−O bonds. Breaking the C−O bond with O binding the Ti atom will produce CH3OH and leave an adsorbed O* on the Ti site. Alternatively, breaking the dangling C−O bond results in an adsorbed methoxy (CH3O*) species on the Ti site and an OH* on Ag. Subsequently, CH3O* can be hydrogenated to produce CH4 and O*. The catalyst will recover its original state after adsorbed O* is protonated to OH* and then to H2O. Based on the free energy profile for the HCOO* path shown in Figure 3, CH3OH formation would be more favorable than CH4 formation, by 0.10 eV. Furthermore, CH3OH formation requires fewer proton reduction and hydrogenation steps, and therefore, should be kinetically favorable.

Figure 6. Catalytic cycle from H2COOH* to produce CH3OH and CH4. Bond lengths are in Å. Ag atoms (blue), Ti atoms (green), O atoms (red) and C atoms (gray).


16

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


3.2.3 Reaction network of CO2 reduction on (TiO2)3/Ag(110) To better understand CO2 reduction mechanism, we constructed a network involving all C1 intermediates in CO2 reduction (Scheme 1) on (TiO2)3/Ag(110). We also highlighted the dominating HCOO* pathway using red arrows. The COOH* pathway was included for completeness.

Scheme 1. Reaction network for CO2 reduction on (TiO2)3/Ag(110). The dominant pathway and key intermediates are highlighted with red arrows and in bold, respectively.

As shown in Scheme 1, CH3OH and CH4 can be produced through either the HCOO* or COOH* intermediate on (TiO2)3/Ag(110). We showed that the HCOO* pathway is the thermodynamically dominating pathway, with H2COOH* being the key intermediate. Further hydrogenation of H2COOH* breaks one of the C−O bonds, resulting in CH3OH and O* or CH3O* and OH*, respectively. The (TiO2)3/Ag(110) catalyst also enables hydrogenation of COOH* to HCOOH* and makes CO* formation energetically unfavorable. HCOOH* formation and its further hydrogenation to H2COOH* merges the COOH* pathway to the HCOO* pathway. We would point out that CH3OH competes with CH4 as the product of the reduction following the formation of CH3O*. Several studies have shown that CH3O* is the key intermediate and determines the selectivity toward CH4 and CH3OH, depending on the oxygen affinity of the


17


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

catalytic sites.50−52 In fact, O binding energy has been used as a descriptor for selectivity between CH3OH and CH4 on metal and metal oxide-based electrocatalysts, and CH4 would be the preferred product on the catalysts exhibiting strong binding toward O.9,33,52 Bhowmik and coworkers reported that the CH3O* intermediate formed on Ti-doped RuO2 prefers to be hydrogenated to CH4 and leaves O* on the surface due to the strong binding ability of Ti toward O.33 On (TiO2)3/Ag(110), the Ti site binds the O end of the CH3O* species with its methyl group (−CH3) pointing away from the metal surface. Hydrogenating at the C position of CH3O* will result in breaking the C−O bond and forming CH4. The high free energy barrier of 1.33 eV to protonate CH3O* to CH3OH makes C−O bond breaking and CH4 formation feasible. Consequently, CH3O* on (TiO2)3/Ag(110) is likely to result in CH4. To sustain the catalytic cycle, the O* species has to be cleared through consecutive protonation to OH* and further to H2O. With a free energy barrier of 1.00 eV, our result indicates that hydrogenation of OH* to H2O is highly endergonic, making the OH* removal the potentiallimiting step. Furthermore, a catalyst such as (TiO2)3/Ag(110) may have to be regenerated/activated to maintain the reactivity and selectivity in practice. Metals with a moderate oxygen affinity are expected to favor methanol over methane following CH3O* formation and also allow more efficient OH removal. 4. CONCLUSIONS We constructed (TiO2)3/Ag(110) as a model of the metal supported metal oxide catalysts for electrochemical reduction of CO2. Our results show that the introduction of metal oxides to a metal electrode provides the interfacial active sites for adsorption and activation of CO2 and facilitates further reduction. On the (TiO2)3/Ag(110) catalyst, Ag acts as an electron-donor and channels the electron to reduce proton and hydrogenating CO2. This catalyst enables CO2


18

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


reduction by facilitating the electron transfer and opens up new reaction channels not available on the pure metal electrode. The dominating pathway on the (TiO2)3/Ag(110) catalyst to produce CH3OH is through H2COOH*, a key intermediate along HCOO* pathway, with the potentiallimiting step being the formation of H2COOH* (0.47 eV). Further hydrogenation to CH4 through protonating H2COOH* and followed by hydrogenation to CH3O* is feasible. The CO* formation from COOH* becomes less favorable on (TiO2)3/Ag(110). The analysis also revealed that all intermediates were anchored to the catalyst through O due to the high affinity of Ti. The strong binding of O-containing species makes OH* removal challenging. In summary, the present study demonstrates that a combination of metal and metal oxide provides new possible catalysts for electrochemical CO2 reduction. Moreover, product selectivity to CH3OH or CH4 can be optimized by tuning the oxygen affinity of the metal in the metal oxide. However, a quantitative estimation of the conversion and selectivity requires extensive knowledge on kinetics as well as mass transport involved in solution chemistry.

ASSOCIATED CONTENT Supporting Information Details of CO2 adsorption structures and DOS analysis, free energy diagrams for CO2 reduction on Ag(110), Bader charges of the Ti atoms in bulk TiO2, Ti2O3 and (TiO2)3/Ag(110) with different intermediates following CO2 reduction, calculated electronic energies and corresponding free energies of adsorbed intermediates on (TiO2)3/Ag(110) at 298.15 K. (PDF)


19


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

ACKNOWLEDGMENT We acknowledge the support of the National Key Research and Development Program of China (2016YFB0600900) and the National Natural Sciences Foundation of China (Grant No. 21206117). We also acknowledge the High Performance Computing Center of Tianjin University and National Supercomputing Center in Shenzhen for providing services to computing resources. QG acknowledges the support of NSF-CBET program (Award no. CBET1438440)

REFERENCES (1) Trends in Atmospheric Carbon Dioxide. ESRL Global Monitoring Division. http://www.esrl.noaa.gov/gmd/ccgg/trends/index.html (accessed Jan 20, 2017). (2) Whipple, D. T.; Kenis, P. J. A. Prospects of CO2 Utilization via Direct Heterogeneous Electrochemical Reduction. J. Phys. Chem. Lett. 2010, 1, 3451−3458. (3) Kondratenko, E. V.; Mul, G.; Baltrusaitis, J.; Larrazábal, G. O.; Pérez-Ramírez, J. Status and Perspectives of CO2 Conversion into Fuels and Chemicals by Catalytic, Photocatalytic and Electrocatalytic Processes. Energy Environ. Sci. 2013, 6, 3112−3135. (4) Thompson, T. L.; Yates, J. T., Jr. Surface Science Studies of the Photoactivation of TiO2 — New Photochemical Processes. Chem. Rev. 2006, 106, 4428−4453. (5) Habisreutinger, S. N.; Schmidt-Mende, L.; Stolarczyk, J. K. Photocatalytic Reduction of CO2 on TiO2 and Other Semiconductors. Angew. Chem. Int. Ed. 2013, 52, 7372−7408. (6) Wang, W.; Wang, S.; Ma, X.; Gong, J. Recent Advances in Catalytic Hydrogenation of Carbon Dioxide. Chem. Soc. Rev. 2011, 40, 3703−3727.


20

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


(7) 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. (8) Zhu, S.; Shao, M. Surface Structure and Composition Effects on Electrochemical Reduction of Carbon Dioxide. J. Solid State Electrochem. 2015, 20, 861−873. (9) 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. (10) 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. (11) Li, Y.; Chan, S. H.; Sun, Q. Heterogeneous Catalytic Conversion of CO2: A Comprehensive Theoretical Review. Nanoscale 2015, 7, 8663−8683. (12) Koppenol, W. H.; Rush, J. D. Reduction Potential of the Carbon Dioxide/Carbon Dioxide Radical Anion: A Comparison with other C1 Radicals. J. Phys. Chem. 1987, 91, 4429−4430. (13) Hori, Y.; Wakebe, H.; Tsukamoto, T.; Koga, O. Electrocatalytic Process of CO Selectivity in Electrochemical Reduction of CO2 at Metal Electrodes in Aqueous Media. Electrochimica Acta 1994, 39, 1833−1839. (14) 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. (15) Hori, Y. Modern Aspects of Electrochemistry; Springer: New York, 2008. (16) Hatsukade, T.; Kuhl, K. P.; Cave, E. R.; Abram, D. N.; Jaramillo, T. F. Insights into the Electrocatalytic Reduction of CO2 on Metallic Silver Surfaces. Phys. Chem. Chem. Phys. 2014, 16, 13814−13819.


21


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

(17) 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. (18) Kumar, B.; Brian, J. P.; Atla, V.; Kumari, S.; Bertram, K. A.; White, R. T.; Spurgeon, J. M. New Trends in the Development of Heterogeneous Catalysts for Electrochemical CO2 Reduction. Catal. Today 2016, 270, 19−30. (19) Lu, Q.; Jiao, F. Electrochemical CO2 Reduction: Electrocatalyst, Reaction Mechanism, and Process Engineering. Nano Energy 2016, 29, 439−456. (20) Zhu, D. D.; Liu, J. L.; Qiao, S. Z. Recent Advances in Inorganic Heterogeneous Electrocatalysts for Reduction of Carbon Dioxide. Adv. Mater. 2016, 28, 3423−3452. (21) Li, Y.; Sun, Q. Recent Advances in Breaking Scaling Relations for Effective Electrochemical Conversion of CO2. Adv. Energy Mater. 2016, 6, 1600463−1600481. (22) Zhu, G.; Li, Y.; Zhu, H.; Su, H.; Chan, S. H.; Sun, Q. Curvature-Dependent Selectivity of CO2 Electrocatalytic Reduction on Cobalt Porphyrin Nanotubes. ACS Catal. 2016, 6, 6294−6301. (23) Li, Y.; Su, H.; Chan, S. H.; Sun, Q. CO2 Electroreduction Performance of Transition Metal Dimers Supported on Graphene: A Theoretical Study. ACS Catal. 2015, 5, 6658−6664. (24) Ma, S.; Lan, Y.; Perez, G. M.; Moniri, S.; Kenis, P. J. Silver Supported on Titania as an Active Catalyst for Electrochemical Carbon Dioxide Reduction. ChemSusChem 2014, 7, 866−874. (25) He, H.; Zapol, P.; Curtiss, L. A. A Theoretical Study of CO2 Anions on Anatase (101) Surface. J. Phys. Chem. C 2010, 114, 21474−21481. (26) Bandi, A. Electrochemical Reduction of Carbon Dioxide on Conductive Metallic Oxides. J. Electrochem. Soc. 1990, 137, 2157−2160.


22

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


(27) Le, M.; Ren, M.; Zhang, Z.; Sprunger, P. T.; Kurtz, R. L.; Flake, J. C. Electrochemical Reduction of CO2 to CH3OH at Copper Oxide Surfaces. J. Electrochem. Soc. 2011, 158, E45−E49. (28) Ramesha, G. K.; Brennecke, J. F.; Kamat, P. V. Origin of Catalytic Effect in the Reduction of CO2 at Nanostructured TiO2 Films. ACS Catal. 2014, 4, 3249−3254. (29) Kim, D.; Lee, S.; Ocon, J. D.; Jeong, B.; Lee, J. K.; Lee, J. Insights into an Autonomously Formed Oxygen-Evacuated Cu2O Electrode for the Selective Production of C2H4 from CO2. Phys. Chem. Chem. Phys. 2015, 17, 824−830. (30) Ullah, N.; Ali, I.; Jansen, M.; Omanovic, S. Electrochemical Reduction of CO2 in an Aqueous Electrolyte Employing an Iridium/Ruthenium-Oxide Electrode. Can. J. Chem. Eng. 2015, 93, 55−62. (31) Karamad, M.; Hansen, H. A.; Rossmeisl, J.; Nørskov, J. K. Mechanistic Pathway in the Electrochemical Reduction of CO2 on RuO2. ACS Catal. 2015, 5, 4075−4081. (32) Gao, S.; Lin, Y.; Jiao, X.; Sun, Y.; Luo, Q.; Zhang, W.; Li, D.; Yang, J.; Xie, Y. Partially Oxidized Atomic Cobalt Layers for Carbon Dioxide Electroreduction to Liquid Fuel. Nature 2016, 529, 68−71. (33) Bhowmik, A.; Vegge, T.; Hansen, H. A. Descriptors and Thermodynamic Limitations of Electrocatalytic Carbon Dioxide Reduction on Rutile Oxide Surfaces. ChemSusChem 2016, 9, 3230−3243. (34) 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


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

(35) Cui, C.; Han, J.; Zhu, X.; Liu, X.; Wang, H.; Mei, D.; Ge, Q. Promotional Effect of Surface Hydroxyls on Electrochemical Reduction of CO2 over SnOx/Sn Electrode. J. Catal. 2016, 343, 257−265. (36) Graciani, J.; Mudiyanselage, K.; Xu, F.; Baber, A. E.; Evans, J.; Senanayake, S. D.; Stacchiola, D. J.; Liu, P.; Hrbek, J.; Fernandez Sanz, J., et al. Highly Active Copper-Ceria and Copper-Ceria-Titania Catalysts for Methanol Synthesis from CO2. Science 2014, 345, 546−550. (37) Kattel, S.; Yan, B.; Yang, Y.; Chen, J. G.; Liu, P. Optimizing Binding Energies of Key Intermediates for CO2 Hydrogenation to Methanol over Oxide-Supported Copper. J. Am. Chem. Soc. 2016, 138, 12440−12450. (38) Kresse, G.; Hafner, J. Ab initio Molecular Dynamics for Liquid Metals. Phys. Rev. B 1993, 47, 558−561. (39) Kresse, G.; Furthmüller, J. Efficiency of Ab-initio Total Energy Calculations for Metals and Semiconductors Using a Plane-Wave Basis Set. Comp. Mater. Sci 1996, 6, 15−50. (40) Kresse, G.; Joubert, D. From Ultrasoft Pseudopotentials to the Projector Augmented-Wave Method. Phys. Rev. B 1999, 59, 1758−1775. (41) Kresse, G.; Furthmüller, J. Efficient Iterative Schemes for Ab initio Total-Energy Calculations Using a Plane-Wave Basis Set. Phys. Rev. B 1996, 54, 11169−11186. (42) Cui, C.; Wang, H.; Zhu, X.; Han, J.; Ge, Q. A DFT Study of CO2 Electrochemical Reduction on Pb(211) and Sn(112). Sci. China Chem. 2015, 58, 607−613. (43) Zhao, Y.; Cui, C.; Han, J.; Wang, H.; Zhu, X.; Ge, Q. Direct C−C Coupling of CO2 and the Methyl Group from CH4 Activation through Facile Insertion of CO2 into Zn−CH3 σ-Bond. J. Am. Chem. Soc. 2016, 138, 10191−10198.


24

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


(44) Li, G.; Han, J.; Wang, H.; Zhu, X.; Ge, Q. Role of Dissociation of Phenol in Its Selective Hydrogenation on Pt(111) and Pd(111). ACS Catal. 2015, 5, 2009−2016. (45) Nørskov, J. K.; Rossmeisl, J.; Logadottir, A.; Lindqvist, L.; Kitchin, J. R.; Bligaard, T.; Jónsson, H. Origin of the Overpotential for Oxygen Reduction at a Fuel-Cell Cathode. J. Phys. Chem. B 2004, 108, 17886−17892. (46) Cramer, C. J. Essentials of Computational Chemistry: Theories and Models. 2nd ed.; John Wiley & Sons: Chichester, England, 2004. (47) Sorescu, D. C.; Al-Saidi, W. A.; Jordan, K. D. CO2 Adsorption on TiO2(101) Anatase: A Dispersion-Corrected Density Functional Theory Study. J. Chem. Phys. 2011, 135, 124701−124717. (48) Yin, W. J.; Krack, M.; Wen, B.; Ma, S. Y.; Liu, L. M. CO2 Capture and Conversion on Rutile TiO2(110) in the Water Environment: Insight by First-Principles Calculations. J. Phys. Chem. Lett. 2015, 6, 2538−2545. (49) Rosen, J.; Hutchings, G. S.; Lu, Q.; Rivera, S.; Zhou, Y.; Vlachos, D. G.; Jiao, F. Mechanistic Insights into the Electrochemical Reduction of CO2 to CO on Nanostructured Ag Surfaces. ACS Catal. 2015, 5, 4293−4299. (50) Peterson, A. A.; Abild-Pedersen, F.; Studt, F.; Rossmeisl, J.; Nørskov, J. K. How Copper Catalyzes the Electroreduction of Carbon Dioxide into Hydrocarbon Fuels. Energy Environ. Sci. 2010, 1311−1315. (51) 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, 2459−2462.


25


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

(52) Back, S.; Kim, H.; Jung, Y. Selective Heterogeneous CO2 Electroreduction to Methanol. ACS Catal. 2015, 5, 965−971.


26

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


TOC Graphics


27

Titania-Modified Silver Electrocatalyst for Selective CO2 Reduction to

Recommend Documents