Patent Application
20240191365
Hydrogen peroxide (H2O2) is a powerful and green oxidant with diverse applications in chemical manufacturing, wastewater treatment, and the paper and pulp industry. The COVID-19 pandemic has also contributed to the recent growth of the global H2O2 market for use in disinfection. The traditional industrial production of H2O2 that proceeds chemically through the anthraquinone process consumes H2 gas and is energy-intensive. It produces up to 70 wt % concentrated H2O2 at centralized plants and requires hazardous transportation to the point-of-use. Decentralized electrosynthesis of H2O2 via the two-electron oxygen reduction reaction (2e− ORR, O2+2 H++2 e− →H2O2) offers a more sustainable route because it can be driven by increasingly affordable renewable electricity, eliminate the need for H2 gas, and produce dilute H2O2 directly at the point-of-use, which is advantageous for distributed applications such as water treatment that only requires dilute (<0.1 wt %) H2O2. The key challenge is to develop robust electrocatalysts with high activity, selectivity, and stability for the desired 2e− reduction to H2O2 (vs. the competing 4e− reduction to water).
Electrochemical valorization of surplus biomass-derived feedstocks, such as glycerol, into high-value chemicals offers a sustainable route for utilizing biomass resources and decarbonizing chemical manufacturing. Glycerol, for example, is typically valorized solely via anodic oxidation, with lower-value products such as hydrogen gas or water generated at cathode. However, there is a need to improve biomass utilization by valorizing biomass resources at both the anode and cathode.
Disclosed herein are electrochemical cells and methods for using the same. One aspect of the invention provides for an electrochemical cell comprising an acidic or neutral electrolyte and a cathode immersed in the electrolyte, wherein the cathode comprises a two-electron oxygen reduction reaction (2e− ORR) electrocatalyst composed of a metal chalcogenide, wherein the metal is Ni or Pd. Another aspect of the technology provides for production of hydrogen peroxide or preparing an oxidation produce of a biomass-derived feedstock with the electrochemical cell.
Another aspect of the technology provides for an electrochemical cell comprising an acidic catholyte, a cathode immersed in the catholyte, an acidic anolyte, and an anode immersed in the anolyte, wherein the cathode comprises a two-electron oxygen reduction reaction (2e− ORR) electrocatalyst composed a metal chalcogenide and the catholyte comprises oxygen (O2), hydrogen peroxide (H2O2), hydroxyl radical (·OH), a regenerable metal ion, and a biomass-derived feedstock, and wherein the anolyte comprises the biomass-derived feedstock. Another aspect of the technology provides for production of hydrogen peroxide or preparing an oxidation product of a biomass-derived feedstock with the electrochemical cell.
Non-limiting embodiments of the present invention will be described by way of example with reference to the accompanying figures, which are schematic and are not intended to be drawn to scale. In the figures, each identical or nearly identical component illustrated is typically represented by a single numeral. For purposes of clarity, not every component is labeled in every figure, nor is every component of each embodiment of the invention shown where illustration is not necessary to allow those of ordinary skill in the art to understand the invention.
Step 2 (blue shaded region): In O2-saturated electrolyte solution, the catalyst-coated disk electrode was linearly swept from 0.75 V to −0.025 vs. RHE (without iR-correction) at the scan rate of 50 mV s−1 and a constant rotation rate (400, 625, 900, 1225, 1600, or 2025 rpm) to drive the acidic 2e− ORR, meanwhile holding the Pt ring electrode at 1.3 V vs. RHE to detect the H2O2 production. The rotation rate was sequentially changed between scans.
Overall procedure: Since holding the Pt ring electrode at 1.3 V vs. RHE for an extended period of time in Step 2 would result in the formation of the surface PtOxS3, S4, the Pt ring electrode was periodically cleaned during the RRDE stability tests using the protocol described in Step 1. Thus, the RRDE stability tests were performed by alternating between Step 1 and Step 2, leading to a total of 255 linear sweep voltammetry scans on the disk electrode over the entire course of ˜-4.0 h. The catalyst stability was monitored by tracking the disk potential required at a certain disk current density (jdisk) or peroxide current density (jperoxide) of 0.5 mA cm−2disk at 2025 rpm (
Disclosed herein is an electrochemical cell and a method of using the same. The present technology allows for harnessing electricity (which could be generated from solar and wind) for electrochemical synthesis of H2O2 and high-value chemicals from biomass feedstocks, which offers a sustainable alternative to conventional centralized chemical manufacturing. For example, glycerol is a byproduct of biodiesel production and has become a surplus biomass-derived chemical. Oxidative upgrading of glycerol is very attractive, because all C3 and C2 oxidation products have higher economic values than glycerol. Compared to thermal oxidation that requires high temperature and oxygen pressure, electrochemical oxidation poses several advantages including near-ambient operation, less reagent waste, and distributed small-scale production.
Electrochemical oxidation of glycerol typically occurs at catalytic anodes made of noble metals or earth-abundant electrocatalysts, which is paired with either four-electron oxygen reduction reaction (4e− ORR) in a galvanic cell or hydrogen evolution reaction (HER) in an electrolytic cell (
An electrochemical cell is defined as a device that may produce an electrical current from a chemical reaction and/or use electrical energy to drive a chemical reaction. The electrochemical cell here includes a cathode immersed in an electrolyte and an anode immersed in an acidic electrolyte. In some embodiments of the disclosed technology, the electrochemical cell may contain two or three electrodes. The two-electrode configuration may contain a working electrode and a counter electrode. The three-electrode configuration may contain a working electrode, a reference electrode, and a counter electrode.
In one aspect, the electrochemical cell includes an electrolyte. An electrolyte is defined as a substance that conducts electric current. The catholyte is defined as an electrolyte used in the cathodic compartment of the electrochemical cell. An anolyte is defined as an electrolyte used in the anodic compartment of the electrochemical cell. The electrolyte, catholyte, and anolyte may be the same material or different material.
In one aspect of the technology, the electrochemical cell contains a cathode with an electrocatalyst that can perform the two-electron oxygen reduction reaction (2e− ORR). The oxygen reduction reaction (ORR) is the reduction half-reaction whereby oxygen is reduced to water via the four-electron pathway or hydrogen peroxide via the two-electron pathway. Hydrogen peroxide (H2O2) is a useful oxidant for a range of applications including the pulp and paper industry, chemical manufacturing, wastewater treatment, healthcare disinfection, and biomass valorization. Other chemical methods of hydrogen peroxide production, including via the anthraquinone process, are energy-intensive and unsafe as it produces concentrated H2O2 at centralized plants and requires hazardous transportation to end-users. Alternatively, decentralized electrosynthesis of H2O2 via 2e− ORR offers a more sustainable route because it can be driven by increasingly affordable renewable electricity, eliminate the need for H2 gas, and produce dilute H2O2 directly at the point of use, which is advantageous for distributed applications. A key challenge is to develop robust electrocatalysts with high activity, selectivity, and stability for the desired 2e (vs. the competing 4e−) ORR pathway.
The thermodynamics of 2e− ORR (O2+2 H++2 e−→H2O2, standard equilibrium potential Eo=0.69 V vs. reversible hydrogen electrode, RHE) and 4e− ORR (O2+4 H++4e−→2 H2O, Eo=1.23 V vs. RHE) are often described by the volcano relations between the thermodynamic limiting potential (UL) and the energetics of key reaction intermediates. 2e− ORR proceeds via the adsorption of OOH* (O2+*+H++e−→OOH*, where * is an unoccupied surface binding site) followed by its desorption to form H2O2 (OOH*+H++e−→H2O2+*); 4e− ORR occurs via the O—O bond cleavage processes (thermal cleavage: O2+2*→2 O*, and OOH*+*→O*+OH*; electrochemical reductive elimination: OOH*+H++e−→O*+H2O). The key intermediates of 2e− ORR (OOH*) and 4e− ORR (OH*) follow a linear scaling relationship (typically ΔGOOH*=ΔGOH*+3.2 eV), resulting in the 2e− and 4e− ORR volcanos. The 2e− ORR activity, determined by the OOH* adsorption energy (ΔGOOH*), is maximized at the peak of 2e− ORR volcano. Moving leftwards from 2e− ORR volcano peak, the catalyst surface binds OOH* (and OH*) more strongly, and UL of 4e− ORR is always more positive than that of 2e− ORR, indicating the 4e pathway will dominate because there is a greater driving force to form H2O than H2O2. To the right of 2e− ORR volcano peak, UL of the 2e− and 4e− pathways overlap and moving rightwards will increase the selectivity (but lowering the activity) for 2e− ORR because the formation of OH* (and OOH*) becomes more difficult. Besides electronic effects described above, the 2e− ORR selectivity can also be improved by controlling geometric (or ensemble) effects by rearranging catalyst surface atoms to change adsorption sites of reaction intermediates, so that O* can be destabilized relative to OOH*, deviating from the conventional scaling relationship.
Several classes of selective 2e− ORR catalysts, including noble metal alloys, carbon nanomaterials, single-atom catalysts, and metal compounds, have been studied for H2O2 electrosynthesis under different pH conditions. Among these reports, 2e− ORR catalysts in alkaline solution have been mostly extensively studied, but H2O2 is unstable in alkaline solution. In contrast, the less studied acidic and neutral conditions are attractive for several reasons besides the chemical stability of H2O2. Acidic H2O2 electrosynthesis can proceed in the technologically mature proton exchange membrane (PEM) devices. On-site water disinfection and environmental treatment can also benefit from acidic H2O2 electrosynthesis because the electro-Fenton process operates at the optimum pH of about 3 to convert the produced H2O2 into the more oxidizing hydroxyl radical (·OH) for the removal of persistent bacteria and organic pollutants. For direct applications, neutral solutions can avoid the need for neutralization. However, high-performance yet cost-effective 2e− ORR electrocatalysts in acidic and neutral solutions are still being developed.
One class of selective 2e− ORR electrocatalysts are metal chalcogenides. In one embodiment of the disclosed technology, the cathode includes an electrocatalyst composed of a metal chalcogenide. A chalcogenide is defined as a chemical compound with at least one chalcogen anion (any anionic form of elements in Group 16 of the periodic table) and at least one more electropositive element. In one embodiment, the chalcogen anion may be a sulfide, selenide, telluride, polonide, or, less commonly, oxide, or any combination thereof.
In one embodiment the electropositive element of the metal chalcogenide is at least one metal. In one embodiment the metal may be earth-abundant or noble metals, including transition metals, post-transition metals, or any combination thereof. In some embodiments, the metal may be Li, Be, Na, Mg, Al, K, Ca, Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn, Ga, Rb, Sr, Y, Zr, Nb, Mo, Tc, Ru, Rh, Pd, Ag, Cd, In, Sn, Cs, Ba, Hf, Ta, W, Re, Os, Ir, Pt, Au, Hg, Tl, Pb, Bi, Po, Fr, Ra. In some embodiments the metal chalcogenide may be cubic pyrite-type c-CoS2, c-CoSe2, c-NiSe2, orthorhombic marcasite-type o-CoSe2, layered PdSe2, layered MoSe2, MoS2, MoSe2, WS2, WSe2, or any combination thereof. In other embodiments, the metal chalcogenide is a spinel such as CuCO2-xNixS4.
Another aspect of 2e− ORR metal chalcogenide electrocatalysts is the preferential binding of O* and OH* products to either the chalcogenide or the metal. Previously studied metal chalcogenide electrocatalysts, including CoS2 and CoSe2, demonstrate preferential binding of O* and OH* to the sulfide and selenide. This has implications for electrocatalyst stability. Surface oxidation is more likely in electrocatalysts where O* and OH* more strongly bind to the chalcogenide due to the formation of highly soluble anions, such as SO42−, and is followed by metal leaching and electrocatalyst degradation. One embodiment of the disclosed technology includes the preferential binding of O* and OH* to the earth abundant metal of the metal chalcogenide. In one embodiment, the electrocatalyst demonstrates preferential binding of O* and OH* to Ni.
In another aspect of the disclosed technology, the electrochemical cell may additionally contain reagents for electrosynthesis purposes. Generally, this involves reagents needed to generate hydroxyl radicals (·OH) via the advanced oxidation processes, including hydrogen peroxide (H2O2), and a regenerable metal ion. In one embodiment, the advanced oxidation may be via the electro-Fenton process, wherein the regenerable metal ion is Fe2+. In other embodiments, the advanced oxidation may be via a Fenton-like process wherein the regenerable metal ion is Co, Cu, or Mn. In some embodiments a regenerable metal ion is not required for the production of the hydroxyl radical. For reasons discussed above, the pH of the electrochemical cell is acidic. In some embodiments, the electrochemical cell has a pH no lower than 0 and no higher than 7. In one embodiment, the electrochemical cell has a pH no lower than 0 and no higher than 4. In another embodiment, the electrochemical cell has a pH between 4.0 and 7.0. In one embodiment where the regenerable metal ion is Fe2+, the electrochemical cell has a pH no lower than 0 and no higher than 4. In some embodiments the pH of the electrochemical cell is adjusted by the addition of a mineral acid. In some embodiments the mineral acid is sulfuric acid, hydrochloric acid, or perchloric acid. In one embodiment where the regenerable metal ion is Fe2+, the pH of the electrochemical cell may be about 3 to prevent the precipitation of iron oxides in more basic conditions. Other regenerable metals may have other ideal acidic, basic, or neutral conditions of the electrochemical cell.
In one embodiment, the electrochemical cell electrolyte further includes a biomass-derived feedstock. Biomass-derived feedstock is defined as sustainable feedstock derived from biomass for chemicals and energy products. In some embodiments, biomass-derived feedstock includes forest product wastes, agricultural residues, organic fractions of municipal solid wastes, paper, cardboard, plastic, food waste, green waste, and other waste. In other embodiments, the biomass-derived feedstock may be polyols such as glycerol, hydroxymethylfurfural (HMF), glucose, fructose, and other monosaccharide and disaccharide (sugar) molecules, lignin, hemicellulose, cellulose, or other oxygen-containing compounds including compounds with ether, alcohol, aldehyde, and ester functional groups.
In some embodiments, the oxidation products of the electrochemical half cell starting from glycerol include C3, C2, and C1 products. In some embodiments the oxidation products include aldehyde and ketone functional groups. In some embodiments the oxidation product may include dihydroxyacetone (DHA), glyceraldehyde (GLAD), glyceric acid (GLA), hydroxypyruvic acid (HPA), glycolaldehyde (GAD), glycolic acid (GA), glyoxylic acid (GLOA), formic acid (FA), or any combination thereof.
Advanced oxidation products are typically only produced at one electrode of the electrochemical cell. In one embodiment, the anode produces oxidation products. One aspect of the disclosed technology is an electrochemical cell wherein electrochemical reactions at both the cathode and anode both produce oxidation products, sometimes referred to as linear paired electrochemical cell.
In another aspect of the technology, the electrochemical cell may contain one or more compartments. In one embodiment, the electrochemical cell is oriented in a two-compartment H-type electrochemical cell (denoted as H-cell), wherein the two compartments are separated by a proton-permeable membrane. In another embodiment, electrochemical flow cell configurations may be used. In electrochemical flow cells, liquid biomass-derived feedstocks, reagents, and electrolytes may be continuously or periodically flowed past electrodes. In one embodiment, the proton permeable membrane may be a Nafion, Fumasep, Fumapem, or Aquivion membrane. In one embodiment the cathode is isolated to one compartment and the anode is isolated to the second compartment. %
In one embodiment, the anode may include platinum, palladium, ruthenium, antimony, bismuth, tin, and bimetallic alloys of these metals, ruthenium oxide, iridium oxide, or boron doped diamond.
In one embodiment the biomass-derived feedstock could be the same in each of the electrochemical cell compartments, representing the linear paired electrochemical process starting with the same feedstock. In another embodiment the biomass-derived feedstock in the cathodic compartment may be different from the biomass-derived feedstock in the anodic compartment, representing the paired electrochemical process in general.
The anodic and cathodic reaction products may be mixed and further purified together but they need not be. Where the anodic and cathodic reaction products are expected to be similar, mixing and purifying the reaction products from both the anodic and cathodic reactions together may simplify processing or reduced costs. In some embodiments, where biomass-derived feedstock is the same in each of the electrochemical cell compartments for the linear paired electrochemical process, the anodic and cathodic reaction products may be mixed and further purified together.
A further aspect of the disclosed technology is a method of preparing oxidation product of biomass-derived feedstock. In one embodiment of the method, the applied potential to the anode is between 0.00 and 1.50 V vs. RHE. In another embodiment, the applied potential to the cathode is between 0.00 and 0.70 V vs. RHE.
Another aspect of the disclosed technology is a method of preparing oxidation product of biomass-derived feedstock with externally applied bias to the two-compartment H-cell. In one embodiment the externally applied bias is less than 0.01 V, 0.02V, 0.03 V, 0.04V, 0.05 V, 0.06 V, 0.07 V, 0.08 V, 0.09 V, 0.10 V, 0.11 V, 0.12 V, 0.13 V, 0.14 V, 0.15 V, 0.16 V, 0.17 V, 0.18 V, 0.19 V, 0.20 V, 0.21 V, 0.22 V, 0.23 V, 0.24 V, 0.25 V, 0.26 V, 0.27 V, 0.28 V, 0.29 V, 0.30 V, 0.31 V, 0.32 V, 0.33 V, 0.34 V, 0.35 V, 0.36 V, 0.37 V, 0.38 V, 0.39 V, 0.40 V, 0.41 V, 0.42 V, 0.43 V, 0.44 V, 0.45 V, 0.46 V, 0.47 V, 0.48 V, 0.49 V, 0.50 V, 0.51 V, 0.52 V, 0.53 V, 0.54 V, 0.55 V, 0.56 V, 0.57 V, 0.58 V, 0.59 V, 0.60 V, 0.61 V, 0.62 V, 0.63 V, 0.64 V, 0.65 V, 0.66 V, 0.67 V, 0.68 V, 0.69 V, 0.70 V, 0.71 V, 0.72 V, 0.73 V, 0.74 V, 0.75 V, 0.76 V, 0.77 V, 0.78 V, 0.79V, 0.80V, 0.81 V, 0.82 V, 0.83 V, 0.84 V, 0.85 V, 0.86 V, 0.87 V, 0.88 V, 0.89 V, 0.90 V, 0.91 V, 0.92 V, 0.93 V, 0.94 V, 0.95 V, 0.76 V, 0.77 V, 0.78 V, 0.79 V, 0.80 V, 0.81 V, 0.82 V, 0.83 V, 0.84 V, 0.85 V, 0.86 V, 0.87 V, 0.88 V, 0.89 V, 0.90 V, 0.91 V, 0.92 V, 0.93 V, 0.94 V, 0.95 V, 0.96 V, 0.97 V, 0.98 V, 0.99 V, 1.00 V, 1.01 V, 1.02V, 1.03 V, 1.04V, 1.05 V, 1.06 V, 1.07 V, 1.08 V, 1.09 V, 1.10 V, 1.11 V, 1.12 V, 1.13 V, 1.14 V, 1.15 V, 1.16 V, 1.17 V, 1.18 V, 1.19 V, 1.20 V, 1.21 V, 1.22 V, 1.23 V, 1.24 V, 1.25 V, 1.26 V, 1.27 V, 1.28 V, 1.29 V, 1.30 V, 1.31 V, 1.32 V, 1.33 V, 1.34 V, 1.35 V, 1.36 V, 1.37 V, 1.38 V, 1.39 V, 1.40 V, 1.41 V, 1.42 V, 1.43 V, 1.44 V, 1.45 V, 1.46 V, 1.47 V, 1.48 V, 1.49 V, or 1.50 V.
By way of example, glycerol valorization in the cathodic half-cell via the electro-Fenton process was studied by applying a fixed potential of 0.60 V vs. RHE to c-NiSe2 grown on carbon fiber paper (denoted as c-NiSe2/CFP) cathode (˜1.06 mgNi cm−2geo, ˜1 cm−2geo) in O2-saturated 0.1 M NaHSO4/Na2SO4 (pH˜2.8) containing ˜50 mM glycerol and 0.5 mM Fe2+, and the cathode current was ˜1.7 mA. Direct oxidation of glycerol in the anodic half-cell was studied by applying a fixed current of 1.7 mA to nanoparticular Pt supported on carbon black (denoted as Pt/C) anode (2 mgPt cm−2geo, 1 cm−2geo) in Ar-saturated 0.05 M H2SO4 (pH˜1.2) containing ˜50 mM glycerol, and the anode potential was ˜0.55 V vs. RHE at pH˜1.2. If c-NiSe2/CFP cathode (in O2-saturated 0.1 M NaHSO4/Na2SO4 containing ˜50 mM glycerol and 0.5 mM Fe2+, pH˜2.8) and Pt/C anode (in Ar-saturated 0.05 M H2SO4 containing ˜50 mM glycerol, pH˜1.2) were coupled together and both operated at 1.7 mA for glycerol valorization in a linear paired electrochemical system (as shown in
Under the ideal assumption of no internal resistance (i.e., no ohmic overpotential) in the linear paired electrochemical system:
This example shows that to deliver a current of ˜1.7 mA for glycerol valorization at both c-NiSe2/CFP cathode and Pt/C anode, the linear paired electrochemical system ideally can operate as an electrolytic cell at an external bias as low as <0.05 V with almost no external energy input needed if the internal resistance is negligible.
The thermodynamic basis of the linear paired glycerol valorization process is provided. Intuitively, the overall process is a controlled partial oxidation of glycerol by oxygen gas to produce mainly C3 and C2 oxidation products in an electrochemical cell, and such oxidation process should be thermodynamically spontaneous.
Considering an ideally simplified linear paired process where both the electro-Fenton process and the direct anodic oxidation process generate a single oxidation product of glyceraldehyde from glycerol, the respective electrochemical reactions at the cathode and the anode should be:
electrochemical reaction at the cathode: O2 (g)+2H++2e−→H2O2 (aq)
electrochemical reaction at the anode: glycerol(aq)→glyceraldehyde(aq)+2H++2e−
The balanced net electrochemical reaction that determines the overall achievable cell potential of this simplified linear paired glycerol-to-glyceraldehyde valorization process should be:
net electrochemical reaction: glycerol(aq)+O2 (g)→glyceraldehyde(aq)+H2O2 (aq)
Because the standard Gibbs free energies of formation (ΔfGo) of these organic compounds are not directly available, we consider the net electrochemical reaction as the combination of the following two reactions:
where ΔGfo[H2O2 (g)] may be sourced from a reference, such as the CRC Handbook of Chemistry and Physics, ΔGsolvo[H2O2 (g)] is the solvation free energy of H2O2 (g) based on the experimental Henry's law constant (kH), and ΔGo for glycerol(aq)→glyceraldehyde(aq)+H2 (g) may be sourced from a theoretical report and may be calculated with quantum chemistry software, such as the Gaussian 03 program, and includes both electrostatic and nonelectrostatic components of the solvation free energies.
As a result, the net electrochemical reaction is thermodynamically spontaneous and results in a galvanic cell with an overall achievable cell potential of 0.49 V:
glycerol(aq)+O2 (g)→glyceraldehyde(aq)+H2O2 (aq)
ΔGo=−134.1 kJ+39.3 kJ=−94.8 kJ
ΔGo/(−n×F)=−94.8 kJ/(−2 mol×96485 C mol−1)=0.49 V
where n is the number of electrons transferred, and F is the Faraday constant.
The linear paired glycerol valorization process is a hybrid of electrochemical and chemical steps. Besides the as-mentioned electrochemical reactions at the cathode and the anode, chemical reaction of glycerol oxidation using H2O2 as the oxygen source takes place in the catholyte solution:
chemical reaction in the catholyte: glycerol(aq)+H2O2 (aq)→glyceraldehyde(aq)+2 H2O(l)
The ΔGo (=−300.8 kJ) of this chemical reaction in the catholyte solution cannot be harnessed electrochemically, and thus the overall achievable cell potential for the linear paired glycerol valorization process is governed by the O2/H2O2 redox couple but not the O2/H2O redox couple. Combining all steps together, the overall cell reaction of this simplified linear paired glycerol-to-glyceraldehyde valorization process should be:
2 glycerol(aq)+O2 (g)→2 glyceraldehyde(aq)+2 H2O(l)
This overall cell reaction is thermodynamically spontaneous (ΔGo=−395.6 kJ), but only part of the ΔGo is harnessed electrochemically to result in a galvanic cell with an overall achievable cell potential of 0.49 V (as discussed above).
Similarly, it holds true for other oxidation products (dihydroxyacetone, glyceric acid, etc.) that the linear paired glycerol valorization process results in a galvanic cell that is thermodynamically spontaneous, which can be derived from the calculated thermochemical data found in ref.
Earlier in this description we estimated based on experimental half-cell studies that, to deliver a current of ˜1.7 mA, our linear paired electrochemical system for glycerol valorization ideally can operate as an electrolytic cell at an external bias as low as <0.05 V if the internal resistance is negligible. In this estimate, the linear paired electrochemical system operates as an electrolytic cell rather than a galvanic cell, because both 2e− ORR at the NiSe2 cathode and glycerol oxidation at the Pt anode require kinetic overpotentials to deliver a catalytic current of ˜1.7 mA. By designing more active cathode and anode electrocatalysts to further reduce the kinetic overpotentials and further optimizing the electrochemical device design, linear paired electrochemical systems for glycerol valorization that need no external bias and no external energy input could be realized.
The linear paired system ideally could operate at an external bias as low as <0.05 V if there was no internal resistance. In
Unless otherwise specified or indicated by context, the terms “a”, “an”, and “the” mean “one or more”. For example, “a molecule” should be interpreted to mean “one or more molecules.”
As used herein, “about”, “approximately,” “substantially,” and “significantly” will be understood by persons of ordinary skill in the art and will vary to some extent on the context in which they are used. If there are uses of the term which are not clear to persons of ordinary skill in the art given the context in which it is used, “about” and “approximately” will mean plus or minus ≤10% of the particular term and “substantially” and “significantly” will mean plus or minus >10% of the particular term.
As used herein, the terms “include” and “including” have the same meaning as the terms “comprise” and “comprising.” The terms “comprise” and “comprising” should be interpreted as being “open” transitional terms that permit the inclusion of additional components further to those components recited in the claims. The terms “consist” and “consisting of” should be interpreted as being “closed” transitional terms that do not permit the inclusion additional components other than the components recited in the claims. The term “consisting essentially of” should be interpreted to be partially closed and allowing the inclusion only of additional components that do not fundamentally alter the nature of the claimed subject matter.
All methods described herein can be performed in any suitable order unless otherwise indicated herein or otherwise clearly contradicted by context. The use of any and all examples, or exemplary language (e.g., “such as”) provided herein, is intended merely to better illuminate the invention and does not pose a limitation on the scope of the invention unless otherwise claimed. No language in the specification should be construed as indicating any non-claimed element as essential to the practice of the invention.
All references, including publications, patent applications, and patents, cited herein are hereby incorporated by reference to the same extent as if each reference were individually and specifically indicated to be incorporated by reference and were set forth in its entirety herein.
Preferred aspects of this invention are described herein, including the best mode known to the inventors for carrying out the invention. Variations of those preferred aspects may become apparent to those of ordinary skill in the art upon reading the foregoing description. The inventors expect a person having ordinary skill in the art to employ such variations as appropriate, and the inventors intend for the invention to be practiced otherwise than as specifically described herein. Accordingly, this invention includes all modifications and equivalents of the subject matter recited in the claims appended hereto as permitted by applicable law. Moreover, any combination of the above-described elements in all possible variations thereof is encompassed by the invention unless otherwise indicated herein or otherwise clearly contradicted by context.
Embodiment 1: An electrochemical cell including an electrolyte and a cathode immersed in the electrolyte,
Embodiment 2: The electrochemical cell of embodiment 1, wherein the metal is Ni.
Embodiment 3: The electrochemical cell of embodiment 2, wherein the metal chalcogenide includes c-NiSe2.
Embodiment 4: The electrochemical cell of any one of embodiments 1-3, wherein the electrolyte has a pH below 4.0, optionally wherein the electrolyte has a pH between 0.0 and 4.0 pH or any pH therebetween.
Embodiment 5: The electrochemical cell of embodiment 1, wherein the metal is Pd.
Embodiment 6: The electrochemical cell of embodiment 5, wherein the metal chalcogenide includes layered PdSe2.
Embodiment 7: The electrochemical cell of any one of embodiments 1 or 5-6, wherein the electrolyte has a pH above 4.0, optionally wherein the electrolyte has a pH between 4.0 and 8.0 pH or any pH therebetween.
Embodiment 8: The electrochemical cell of any one of embodiments 1-7, wherein the electrolyte includes oxygen (O2), hydrogen peroxide (H2O2), hydroxyl radical (·OH), a regenerable metal ion, and a first biomass-derived feedstock.
Embodiment 9: The electrochemical cell of embodiment 8 further including an acidic anolyte and an anode immersed in the anolyte, wherein the anolyte includes a second biomass-derived feedstock.
Embodiment 10: A method for production of hydrogen peroxide, the method including introducing oxygen into the electrochemical cell according to any one of embodiments 1-9 under conditions sufficient for preparing the hydrogen peroxide.
Embodiment 11: The method of embodiment 10, wherein the conditions sufficient for preparing hydrogen peroxide include a potential between 0.00 and 0.70 V vs. RHE applied to the cathode.
Embodiment 12: An electrochemical cell including an acidic catholyte, a cathode immersed in the catholyte, an acidic anolyte, and an anode immersed in the anolyte,
Embodiment 13: The electrochemical cell of embodiment 12, wherein the two-electron oxygen reduction reaction (2e− ORR) electrocatalyst includes a metal chalcogenide.
Embodiment 14: The electrochemical cell of embodiment 13, wherein the metal chalcogenide includes c-NiSe2.
Embodiment 15: The electrochemical cell of embodiment 13, wherein the metal chalcogenide includes layered PdSe2.
Embodiment 16: The electrochemical cell of embodiment 13, wherein the metal chalcogenide includes pyrite or marcasite type CoSe2, pyrite type CoS2, or CuCO2-xNixS4.
Embodiment 17: The electrochemical cell of embodiment 13, wherein the metal chalcogenide includes an earth-abundant metal or a noble metal.
Embodiment 18: The electrochemical cell of embodiment 13, wherein the earth abundant metal includes Ni, Co, Fe, Cu, Mn, or any combination thereof.
Embodiment 19: The electrochemical cell of embodiment 18, wherein the earth abundant metal includes Ni.
Embodiment 20: The electrochemical cell of embodiment 18, wherein the earth abundant metal includes Co.
Embodiment 21: The electrochemical cell of embodiment 17, wherein the noble metal includes Pd.
Embodiment 22: The electrochemical cell of any one of embodiments 12-21, wherein the acidic catholyte and/or the acidic anolyte has a pH between 0.0 and 4.0 pH.
Embodiment 23: The electrochemical cell of any one of embodiments 12-22, wherein the regenerable metal ion is Fe2+.
Embodiment 24: The electrochemical cell of any one of embodiments 12-23, wherein the electrochemical cell includes a semipermeable barrier between the cathode from the anode.
Embodiment 25: The electrochemical cell of any one of embodiments 12-24, wherein the first biomass-derived feedstock or the second biomass-derived feedstock includes glycerol.
Embodiment 26: The electrochemical cell of any one of embodiments 12-25, wherein each of the catholyte and the anolyte further include oxidation products of the first biomass-derived feed stock and the second biomass derived feed stock.
Embodiment 27: The electrochemical cell of embodiment 26, wherein the oxidation products include dihydroxyacetone (DHA), glyceraldehyde (GLAD), glyceric acid (GLA), hydroxypyruvic acid (HPA), glycolaldehyde (GAD), glycolic acid (GA), glyoxylic acid (GLOA), formic acid (FA), or any combination thereof.
Embodiment 28: A method for preparing an oxidation product of a biomass-derived feedstock, the method including introducing the biomass-derived feedstock into the electrochemical cell according to any one of embodiments 1-9 under conditions sufficient for preparing the oxidation product.
Embodiment 29: The method of embodiment 28, wherein a potential is between 0.00 and 0.70 V vs. RHE is applied to the cathode.
Embodiment 30: The method of any one of embodiments 28-29, wherein the biomass-derived feedstock includes glycerol.
Embodiment 31: The method of any one of embodiments 28-30, wherein the electrolyte further includes an oxidation product of the biomass-derived feedstock.
Embodiment 32: The method of embodiment 31, wherein the oxidation product includes dihydroxyacetone (DHA), glyceraldehyde (GLAD), glyceric acid (GLA), hydroxypyruvic acid (HPA), glycolaldehyde (GAD), glycolic acid (GA), glyoxylic acid (GLOA), formic acid (FA), or any combination thereof.
Embodiment 33: A method for preparing an oxidation product of a biomass-derived feedstock, the method including introducing the biomass-derived feedstock into the electrochemical cell according to any one of embodiments 12-24 under conditions sufficient for preparing the oxidation product.
Embodiment 34: The method of embodiment 33, wherein less than 1.0 V of externally applied bias is applied.
Embodiment 35: The method of embodiment 33, wherein less than 0.2 V of externally applied bias is applied.
Embodiment 36: The method of any one of embodiments 33-35, wherein the biomass-derived feedstock includes glycerol.
Embodiment 37: The method of any one of embodiments 33-36, wherein the electrolyte further includes an oxidation product of the biomass-derived feedstock.
Embodiment 38: The method of embodiment 37, wherein the oxidation product includes dihydroxyacetone (DHA), glyceraldehyde (GLAD), glyceric acid (GLA), hydroxypyruvic acid (HPA), glycolaldehyde (GAD), glycolic acid (GA), glyoxylic acid (GLOA), formic acid (FA), or any combination thereof.
Linear paired electrochemical valorization of glycerol requires a cathodic reaction that can generate oxidative species to oxidize glycerol. Hydrogen peroxide (H2O2) is an oxidant (Eo=1.76 V vs. SHE) that can be cathodically generated via the selective 2e− ORR (O2+2 H++2 e−→H2O2), and be further converted into the even more oxidizing hydroxyl radical (·OH, Eo=2.80 V vs. SHE) by the Fe2+-mediated electro-Fenton process in acidic solutions (Fe2++H2O2+H+→Fe3++H2O+·OH) where Fe2+ is regenerated at the H2O2-generating cathode (Fe3++e−→Fe2+) The application of electro-Fenton process has been largely limited to environmental pollutant removal, but chemically generated ·OH from H2O2 has found use in biomass-to-chemical processes such as carbohydrate oxidation and lignin depolymerization.
Here, we present the cathodic valorization of glycerol via the electro-Fenton process, and the further linear pairing with the anodic oxidation to concurrently produce the same glycerol-derived oxidation products at both cathode and anode (
Identifying c-NiSe2 Catalyst for the Electro-Fenton Process
The Fe2+-mediated electro-Fenton process operates at an optimum pH of ˜3 and poses more stringent requirements for catalyst stability than 2e− ORR because OH is more oxidizing than H2O2. Therefore, an electrocatalyst that is not only selective for acidic 2e− ORR but also stable in the presence of strong oxidants such as H2O2 and ·OH is needed. We utilized the calculated bulk Pourbaix diagrams available from the Materials Project to identify promising earth-abundant catalyst candidates with high aqueous electrochemical stability in the pH and potential ranges of interest for acidic 2e− ORR. Similar to cubic pyrite-type CoSe2 (c-CoSe2,
The promise of c-NiSe2 as an active and selective 2e− ORR catalyst is revealed by the calculated free energy diagrams of the ORR energetics on the most thermodynamically stable (100) surface. The 2e− ORR (
The surface stability of c-NiSe2 under aqueous electrochemical environments is evaluated by considering O* and/or OH* adsorbate formation when the surface is in equilibrium with water. Unlike c-CoSe2 where O* and OH* preferentially bind to Se (Se—O*) and Co (Co—OH*), respectively, Ni on c-NiSe2 is the preferential binding site for both O* (Ni—O*) and OH* (Ni—OH*). On a surface unit cell including two metal sites and four Se sites, should O* builds up on the c-NiSe2 surface, a significant O* coverage would have to be reached (which is unlikely because O* binds to Ni endothermically by 0.08 eV at URHEo) before any O* would bind to Se; however, any presence of O* on c-CoSe2 would bind to Se immediately (
Electrocatalytic Properties and Stability of c-NiSe2 for Acidic 2e− ORR
We synthesized nanostructured c-NiSe2 (
The stability of c-NiSe2 (vs. c-CoSe2) catalyst for acidic 2e− ORR was evaluated by long-term RRDE stability tests at various catalyst loadings. The catalyst stability is monitored by tracking the disk potential at a certain disk current density (jdisk) or peroxide current density (jperoxide) (
[1] These avg. ± std. dev. leaching rates come from four RRDE stability tests of c-NiSe2 at different catalyst loadings in each electrolyte, as tabulated in Table 2.
[2] These leaching rates come from the initial H2O2 bulk electrosynthesis run of c-NiSe2/CFP at 0.60 V vs. RHE in each electrolyte, as shown in FIG. 23d.
Bulk Electrosynthesis of H2O2 in Acidic Solution Using c-NiSe2 Cathode
We further performed constant-potential bulk electrosynthesis using integrated electrodes of c-NiSe2 nanosheets directly grown on carbon fiber paper (NiSe2/CFP,
To understand this potential-dependent electrosynthesis of H2O2, we studied the side reaction of H2O2 electroreduction in competition with 2e− ORR on c-NiSe2 catalyst drop-casted on RRDE. In 0.05 M H2SO4, the catalytic onset potential of H2O2 electroreduction on c-NiSe2 coincides with that of 2e− ORR, and the rate of H2O2 electroreduction increases with higher overpotential and H2O2 concentration (
We demonstrated sustained bulk electrosynthesis of H2O2 in O2-saturated 0.05 M H2SO4 at the optimum 0.60 V vs. RHE using one NiSe2/CFP electrode repeatedly for five consecutive runs over 37 hours (
Glycerol Valorization Via the Electro-Fenton Process at c-NiSe2 Cathode
To enable glycerol valorization by the electro-Fenton process, we operated NiSe2/CFP cathode at the fixed potential of 0.60 V vs. RHE in O2-saturated 0.1 M NaHSO4/Na2SO4 buffer (pH˜2.8) containing Fe2+ and glycerol. The balanced equation shows that cathodic glycerol conversion consumes protons (
We further studied the impact of Fe2+ concentration ([Fe2+]) on the glycerol valorization via the electro-Fenton process. The rate of OH formation from the Fenton reaction should increase with higher [Fe2+ ] based on the rate law, but too much Fe2+ would consume the formed OH and decrease the oxidizing power (Fe2+·OH+H+→Fe3++H2O). After a controlled amount of charge is passed through NiSe2/CFP cathode at 0.60 V vs. RHE, high glycerol conversion is achieved when [Fe2+] is 0.5 mM or 1.0 mM, while too little Fe2+ (0.1 mM) results in low glycerol conversion likely due to the slow OH formation (
[1] GLAD = glyceraldehyde; DHA = dihydroxyacetone; GLA = glyceric acid; HPA = hydroxypyruvic acid; GAD = glycolaldehyde; GA = glycolic acid; GLOA = glyoxylic acid; FA = formic acid;
[2] See Examples for details in the Faradaic efficiency calculation.
Pairing the Electro-Fenton Process with Anodic Oxidation for Glycerol Valorization.
To valorize glycerol at both cathode and anode concurrently, anodic glycerol oxidation needs to operate in acidic solution to match the pH requirement of the electro-Fenton process. Therefore, anodic glycerol oxidation was performed in an Ar-saturated H2SO4 solution containing 50 mM glycerol on an anode made by drop-casting commercial Pt/C catalyst on carbon fiber paper. This paired system needs to operate in a two-compartment H-cell rather than in an undivided cell because the O2 needed for the electro-Fenton process can undergo undesirable ORR on the Pt/C anode. Protons are transported through Nafion membrane and stabilize the pH of the catholyte (O2-saturated NaHSO4/Na2SO4 buffer containing 50 mM glycerol and 0.5 mM Fe2+, pH˜2.8) where the electro-Fenton process takes place.
Anodic glycerol oxidation at Pt/C anode in 0.05 M H2SO4 was first evaluated in the half-cell (
We then demonstrated the proof-of-concept linear paired electrochemical valorization of glycerol by feeding glycerol in both cathode and anode compartments of the H-cell where NiSe2/CFP cathode was operated at 0.60 V vs. RHE and Pt/C anode matched the current (
When a higher supporting electrolyte concentration of 0.5 M was applied for both catholyte and anolyte, the paired system operated at a much lower external bias below 0.2 V (
In summary, we demonstrated a linear paired electrochemical process for concurrent glycerol valorization by the electro-Fenton process at a stable and earth-abundant cathode together with direct oxidation at an anode. This process is enabled by the development of highly selective and stable 2e− ORR catalyst for H2O2 production in acidic solution, which is elucidated by calculated free energy diagrams and surface adsorbate analyses and experimentally shown with RRDE and catalyst leaching studies together with sustained electrosynthesis of H2O2. The electro-Fenton process at the cathode at the demonstrated operation conditions leads to efficient cathodic glycerol valorization with a high selectivity toward valuable C3 oxidation products and high glycerol conversion of 55%. The linear paired system achieves similarly high glycerol conversion and product selectivity and can operate at a very small external bias below 0.2 V, which could made into an unbiased system after further optimization in the future. The design principles for stable and selective electrocatalysts for acidic H2O2 production and the electro-Fenton process, and the conceptual strategy of linear pairing the electro-Fenton process with anodic oxidation presented here open up new opportunities for electrochemical valorization of a variety of biomass feedstocks with high atom efficiency and low energy cost.
Spin polarized electronic structure calculations were performed using the Vienna Ab initio Simulation package (VASP) interfaced with the Atomic Simulation Environment (ASE). Projector augmented wave (PAW) pseudopotentials with a cutoff of 450 eV were used to treat core electrons, and the Perdew-Burke-Ernzerhof (PBE) functional was used to treat exchange and correlation. Dispersion was treated using Grimme's D3(ABC) method. To better describe the Co 3d electrons in c-CoSe2, a Hubbard U parameter, Ueff=2.0 eV, was taken from a previous report. A variety of Hubbard U parameters were tested for c-NiSe2, and were found to have little to no effect on the geometries or energies; therefore, no Hubbard U parameter was used for this catalyst. Solvation effects were treated using the continuum solvent method VASPsol. The Brillouin zone was sampled using a (10, 10, 10) and (10, 10, 1) Γ-centered Monkhorst-Pack mesh for bulk and surface calculations, respectively. Lattice constants were determined by fitting to an equation of state (EOS).
For both c-NiSe2 and c-CoSe2, their respective (100) surface exhibits the lowest surface energy compared to other crystal surfaces, and thus is the most thermodynamically stable surface [0.044 vs. 0.064 vs. 0.069 eV Å−2 for c-CoSe2 (100) vs. (110) vs. (111) surface; 0.036 vs. 0.053 vs. 0.058 eV Å−2 for c-NiSe2 (100) vs. (110) vs. (111) surface]. The (100) surfaces of c-NiSe2 and c-CoSe2 were modelled as a 1×1 unit cell slab with two repeats in the z-direction, leading to a total of 8 metal atoms and 16 Se atoms and a vacuum gap of at least 15 Å. The top half of the slabs was allowed to relax while the bottom half was frozen to simulate the bulk. For each ionic configuration, the electronic energy was converged below 10−6 eV. Both the clean slab and adsorbates were allowed to relax until the forces were converged below 0.005 eV Å−2. Transition states were located using the nudged elastic band (NEB) method and were refined using the dimer method. All transition states were confirmed saddle points with one imaginary frequency corresponding to bond breaking.
Binding energies were calculated with respect to O2 (g) and H+(aq) and e−. The energy of H+(aq) and e was calculated using the computational hydrogen electrode (CHE) method, where H+(aq) is assumed to be in thermodynamic equilibrium with H2 (g). The use of the CHE method for our calculation is validated by the fact that the difference in the OOH* binding energy on the c-NiSe2 (100) surface calculated by the CHE method vs. the grand-canonical density functional theory (GC-DFT) method, accounting for the change in surface charge density upon adsorption, is <0.1 eV and can be safely neglected. In order to avoid well-known errors in the DFT treatment of O2 (g), the free energy of O2 (g) was determined by matching the experimental standard equilibrium potential (1.229 V) of the reaction ½ O2 (g)+2 H+(aq)+2 e−→H2O(l). The adsorption of O2, forming O2* from O2 (g), is excluded from our calculation because DFT does not treat O2 (g) accurately, and the estimated free energy difference between O2 (g) and O2* on the c-NiSe2 (100) surface is <0.1 eV and can be safely neglected. The free energies of species were calculated using G=H−T·S, where H is the enthalpy including zero-point energy (ZPE) and thermal corrections, and S is either the total experimental entropy at 298 K and 1 bar (for gas phase species) or calculated under the harmonic approximation taking into account both vibrational contributions and hindered translations/rotations (for surface bound species). The free energy of H2O(l) was calculated using the experimental free energy of formation for H2O(l) and H2O(g). The solvation free energy of H2O2 (aq) was calculated using the experimental Henry's law constant. The calculated standard equilibrium potential (URHEo) of the 2e− ORR reaction O2 (g)+2 H+(aq)+2 e−→H2O2 (aq) is 0.81 V, slightly higher than the experimental standard equilibrium potential of 0.69 V.
Free energies of different surface adsorbate coverages were calculated to predict the most thermodynamically stable surface termination of each catalyst for a given set of potential and pH conditions under the assumption that the surfaces can be approximated in equilibrium with H2O(l). The equilibrated proton-coupled electron transfer reaction for a general surface intermediate can then be written as:
X—OmHn*+(2m−n)(H++e−)⇄X*+mH2O
where X is the surface binding site, m is the number of oxygen atoms, and n is the number of hydrogen atoms. The free energy of this reaction can be written as:
ΔG(U,pH)=GX*+mGH
Using the computational hydrogen electrode (CHE) method (Ge
ΔG(URHE)=GX*+mGH
A 1×1 unit cell slab of the (100) surface of each catalyst that has two metal binding sites and four Se binding sites was used to model intermediate surface coverages as a function of potential. For c-NiSe2, the Ni site is the preferential binding site for both OH* and O*. For c-CoSe2, the Co site is the preferential binding site for OH*, and the Se site is the preferential binding site for O*. A wide variety of surface coverages were examined on various combinations of binding sites. For the sake of clarity, only the most thermodynamically stable surface coverages (in the URHE range of 0 V to 0.95 V) on the most preferential combination of binding sites were shown in
All chemicals were purchased from Sigma-Aldrich and used as received without further purification, unless noted otherwise. Deionized nanopure water (18.2 MΩ·cm) from Thermo Scientific Barnstead water purification systems was used for all experiments.
c-NiSe2 powder sample was prepared by a hydrothermal method. Following a procedure modified from a previous report, nickel hydroxide [Ni(OH)2] precursor was first synthesized by dissolving 451.3 mg of NiSO4·6H2O (Acros Organics, 98+%) in 58.3 mL of water and 8.75 mL of 2 M ammonia aqueous solution (diluted from ammonium hydroxide solution, 28.0-30.0% NH3 basis), and heating the solution at 180° C. for 24 h in a sealed 100-mL Teflon-lined stainless-steel autoclave. The resulting Ni(OH)2 precursor was hydrothermally converted into c-NiSe2 as follows: 4.29 g of NaOH (≥97.0%) and 571 mg of Se powder (≥99.5%) were suspended in 50 mL of water via sonication and heated at 220° C. for 24 h in a sealed 80-mL autoclave; upon cooling to room temperature, 35 mg of Ni(OH)2 precursor was suspended in 10 mL of water and added dropwise into the Se-containing solution under vigorous stirring, and then heated at 180° C. for additional 24 h in the same autoclave. The as-converted c-NiSe2 product was successively washed with water, 1.25 M aqueous solution of Na2S (nonahydrate, ≥98.0%) (to dissolve the elemental Se impurity), and water four times for each washing step, and dried under vacuum at 50° C.
To prepare Ni(OH)2 precursor on carbon fiber paper (Ni(OH)2/CFP), Teflon-coated carbon fiber paper (Fuel Cell Earth, TGP-H-060) was first treated with oxygen plasma at 150 W power for 5 min for each side and annealed in air at 700° C. for 5 min. A 3 cm×6 cm piece of annealed CFP was placed in the solution made of 2.1 mmol of Ni(NO3)2·6H2O (≥97.0%), 4.2 mmol of NH4F (≥98.0%), and 10.5 mmol of urea (99.0-100.5%) in 80 mL of water, and heated at 110° C. for 5 h in a sealed 100-mL autoclave. NiSe2/CFP was prepared by the same hydrothermal selenization method described above, except for using a 1.5 cm×6 cm piece of Ni(OH)2/CFP as the precursor. The as-converted NiSe2/CFP was immersed in 1.25 M aqueous solution of Na2S three times to remove any excess elemental Se, rinsed with water and ethanol, and dried under N2 gas flow. The areal loading of c-NiSe2 grown on CFP was determined by the mass change of CFP after the materials growth. The c-CoSe2 samples were prepared following the published procedures. All catalyst samples were stored in an argon-filled glove box to minimize the exposure to air.
Powder X-ray diffraction (PXRD) patterns were collected on a Bruker D8 ADVANCE powder X-ray diffractometer using Cu Kα radiation. Scanning electron microscopy (SEM) was performed on a Zeiss SUPRA 55VP field emission scanning electron microscope at an accelerating voltage of 1 kV. For SEM imaging, powder samples were drop-casted onto silicon wafer substrates. X-ray photoelectron spectroscopy (XPS) was performed on a Thermo Scientific K-Alpha XPS system with an Al Kα X-ray source. Raman spectroscopy was performed on a Thermo Fisher Scientific DXR3xi Raman Imaging Microscope using a 50 m confocal pinhole aperture and a 532 nm laser source and with a low laser power of 0.1 mW and an exposure time of 1.0 second to avoid sample degradation. For XPS and Raman experiments, powder samples were drop-casted onto graphite substrates, which were made by cutting thin slices of graphite rod (Graphite Store, low wear EDM rod), abrading with 600 grit silicon carbide paper (Allied High Tech Products), and sonicating in water and ethanol until clean. X-ray absorption spectroscopy (XAS) of NiSe2/CFP before and after electrochemical testing was performed in transmission mode at the Advanced Photon Source (APS) Beamline 10-BM-B, and analyzed using ATHENA and ARTEMIS software.
Drop-casted catalysts were prepared on a rotating ring-disk electrode (RRDE-3A, ALS Co., Ltd) made of a glassy carbon disk (with a geometric area of 0.126 cm2disk) surrounded by a Pt ring. The collection efficiency of the bare RRDE was 0.43, determined experimentally using the ferri-/ferrocyanide redox couple. The RRDE was successively polished with 1, 0.3, and 0.05 m alumina suspensions (Allied High Tech Products) on a polishing cloth (Buehler, MicroCloth), thoroughly rinsed with water and methanol, briefly sonicated in methanol for <20 s, and dried under ambient conditions before use. The catalyst inks were prepared by suspending pre-weighed catalyst powders in desired volumes of a 9:1 (v/v) mixture of water and 5 wt % Nafion solution by sonication for 1 h. A fixed volume (10 μL) of catalyst ink was then drop-casted on the disk electrode and dried under ambient conditions at a rotation rate of 700 rpm to form a uniform catalyst film where the Nafion loading was identical (0.4 mgNafion cm−2disk) whereas the catalyst loading was varied (
RRDE measurements were conducted in an undivided three-electrode cell with a graphite rod counter electrode and a Hg/Hg2SO4 (saturated K2SO4) reference electrode (calibrated against a saturated calomel electrode) connected to a Bio-Logic VMP-300 multichannel potentiostat. All potentials were reported versus RHE (E vs. RHE=E vs. SHE+0.059×pH). Prior to RRDE measurements, the electrolyte solution (40-45 mL) of either 0.05 M H2SO4 (pH˜1.2) or 0.1 M NaHSO4/Na2SO4 buffer solution (pH˜2.8) was purged with O2 gas for >10 min, and a blanket of O2 gas was maintained above the electrolyte solution during the measurements. Under O2-saturated condition, the catalyst-coated disk was first conditioned by running cyclic voltammetry (CV) between −0.025 V and 0.75 V vs. RHE at 100 mV s−1 and 1600 rpm for 10 cycles, while holding the Pt ring at 1.3 V vs. RHE. The Pt ring was then conditioned by running CV between 0.05 V and 1.20 V vs. RHE at 100 mV s−1 and 1600 rpm for 10 cycles while holding the disk at 0.75 V vs. RHE to remove the surface PtOx. The 2e− ORR catalytic properties were evaluated by performing linear sweep voltammetry of the catalyst-coated disk from 0.75 to −0.025 V vs. RHE at 50 mV s−1 and various rotation rates, meanwhile holding the Pt ring at 1.3 V vs. RHE. Finally, the background current, double-layer capacitance (Cdl, determined by CV of the disk between −0.025 V and 0.75 V vs. RHE at various scan rates and 0 rpm), and uncompensated resistance (Ru, determined by electrochemical impedance spectroscopy of the disk at 0.75 V vs. RHE) were measured under Ar-saturated conditions. By manually conducting background current and iR corrections, the H2O2 selectivity (PRRDE) is calculated as follows:
where idisk and iring are the respective disk and ring current, and N is the collection efficiency. For the ease of visualizing the H2O2 selectivity from RRDE voltammograms (
where jperoxide is the partial current density for H2O2 production.
The protocols for long-term RRDE stability tests were described in
Bulk Electrosynthesis of H2O2
NiSe2/CFP cathode (vide supra) was used for constant-potential bulk electrosynthesis of H2O2 in O2-saturated 0.05 M H2SO4 (pH˜1.2) or 0.1 M NaHSO4/Na2SO4 (pH˜2.8) solution (4 mL, stirred at 1200 rpm) placed in the cathode compartment of a two-compartment three-electrode H-cell (see
where [H2O2] is the cumulative H2O2 concentration, [Ce4+]before and [Ce4+]after are the [Ce4+] in the stock solution (determined by fitting to the standard curve) before and after mixing with the catholyte aliquot. The cumulative H2O2 yield (nH
where F is the Faraday constant (96485 C mol−1). To monitor the catalyst leaching during H2O2 bulk electrosynthesis, the spent catholytes were filtered with 0.22-μm syringe filters (Restek) and diluted by 15 times with a matrix solution of 0.05 M H2SO4 for ICP-MS analysis (vide supra).
All experiments of glycerol valorization were performed in the two-compartment three-electrode H-cell described above. Half-cell studies of glycerol valorization via the electro-Fenton process at NiSe2 cathode were performed by chronoamperometry with controlled amounts of charge passed at 0.60 V vs. RHE in O2-saturated 0.1 M NaHSO4/Na2SO4 solution (pH˜2.8) containing glycerol (˜50 mM) and Fe2+ (0.1, 0.5, 1.0, or 2.5 mM, prepared from FeSO4·7H2O, ≥99.0%) (see schematic in
where Vi and Vf are the initial and final electrolyte volume, [glycerol]i and [glycerol]f are the initial and final concentration of glycerol, [Cn product]f is the final concentration of Cn product (n=1, 2, 3), all of which are listed in Table 3 and 5. The Faradaic efficiency of all detected aqueous phase organic products at the anode (FEanode) and the cathode (FEcathode) are calculated and estimated, respectively, based on the methods described in below and these calculated FEanode and estimated FEcathode values are also listed in Table 3 and 5.
[1] See FIG. 30a.
[2] See FIG. 30b.
[1] GLAD = glyceraldehyde; DHA = dihydroxyacetone; GLA = glyceric acid; HPA = hydroxypyruvic acid; GAD = glycolaldehyde; GA = glycolic acid; GLOA = glyoxylic acid; FA = formic acid.
Here we provide the details in the Faradaic efficiency (FE) calculation for glycerol valorization at the NiSe2 cathode and the Pt anode.
The FE of all detected aqueous phase organic products (i.e., not including the potential product of CO2, which is not detectable by 1H NMR) from glycerol oxidation at the Pt anode (FEanode) can be calculated using the following equation:
where Vf is the final electrolyte volume, [Cn product]f is the final concentration of Cn product (n=1, 2, 3), x is the theoretical number of electrons transferred for oxidizing 1 molecule of glycerol to 3/n molecule(s) of a specific Cn product at anode (x=2 for glyceraldehyde, dihydroxyacetone, or glycolaldehyde; x=4 for glyceric acid; x=5 for glycolic acid; x=6 for hydroxypyruvic acid; x=8 for glyoxylic acid, or formic acid) based on the balanced anodic half-cell reactions below (see
glycerol→glyceraldehyde+2H++2e−
glycerol→dihydroxyacetone+2H++2e−
glycerol+H2O→glyceric acid+4H++4e−
glycerol+H2O→hydroxypyruvic acid+6H++6e−
glycerol→3/2 glycolaldehyde+2H++2e−
glycerol+3/2H2O→3/2 glycolic acid+5H++5e−
glycerol+3/2H2O→3/2 glyoxylic acid+8H++8e−
glycerol+3H2O→3 formic acid+8H++8e−
The calculated FEanode values for the anodic half-cell studies (63-85%) and the linear paired experiments (63-71%) are listed in Table 3 and 5, respectively. These calculated FEanode values are less than unity due to the oxidation of formic acid into CO2 at the Pt anode. An earlier literature also showed the possible formation of CO2 from glycerol oxidation at Pt anode under similar applied potential and pH conditions.
On the other hand, the estimate of the FE for the cathodic electro-Fenton process is complicated by the multi-step process. We assume that cathodic valorization of glycerol theoretically proceeds via electrogeneration of H2O2 from O2 (O2+2 H++2 e−→H2O2), followed by using H2O2 as the oxygen source that is using the ·OH and Fe3+ pair generated from the Fenton reaction (Fe2++H2O2→Fe3++·OH+OH−) as the oxidants to oxidize glycerol in the solution.
The balanced cathodic half-cell reactions should theoretically be (see
glycerol+O2+2H++2e−→glyceraldehyde+2H2O
glycerol+O2+2H++2e−→dihydroxyacetone+2H2O
glycerol+2O2+4 H++4e−→glyceric acid+3H2O
glycerol+3O2+6 H++6e−→hydroxypyruvic acid+5H2O
glycerol+O2+2H++2e−→3/2 glycolaldehyde+2H2O
glycerol+5/2O2+5 H++5e−→3/2 glycolic acid+7/2 H2O
glycerol+4O2+8 H++8e−→3/2 glyoxylic acid+13/2 H2O
glycerol+4O2+8 H++8e−→3 formic acid+5H2O
The FE of all detected aqueous phase organic products from glycerol valorization at the NiSe2 cathode (FEcathode) can be estimated using the following equation:
where y is the theoretical number of electrons transferred for oxidizing 1 molecule of glycerol to 3/n molecule(s) of a specific Cn product (n=1, 2, 3) at cathode as shown in the balanced cathodic half-cell reactions above (y=2 for glyceraldehyde, dihydroxyacetone, or glycolaldehyde; y=4 for glyceric acid; y=5 for glycolic acid; y=6 for hydroxypyruvic acid; y=8 for glyoxylic acid, or formic acid). Note that this would be a conservative lower bound estimate of the FEcathode, because any non-ideal electro-Fenton process will lead to larger y values and thus larger FEcathode values. The estimated FEcathode values for the cathodic half-cell studies (52-58% when [Fe2+]=0.5 mM) and the linear paired experiments (62-69% when [Fe2+]=0.5 mM) are listed in Table 3 and 5, respectively. These estimated FEcathode values are less than unity due to the oxidation of formic acid into CO2 at the NiSe2 cathode.
Synthesis and Characterization of PdSe2 Nanoplates. We synthesized layered PdSe2 (
Evaluation of the 2e− ORR Activity, Selectivity and Operational Stability of PdSe2. To evaluate the 2e− ORR activity, selectivity, and stability of PdSe2, electrochemical tests using a rotating ring disk electrode (RRDE) in acidic (0.05 M H2SO4, pH=1.2) and neutral (0.05 M NaPi buffer=0.025 M Na2HPO4/0.025 M NaH2PO4, pH=6.5) conditions (
Bulk Electrosynthesis of H2O2 on PdSe2. To evaluate the practical performance of PdSe2 for electrosynthesis of H2O2, Inventors prepared a gas diffusion electrode (GDE) loaded with PdSe2 catalyst (
Based on the CV, Inventors performed an extended electrosynthesis test at a fixed current of −30 mA (
The drops in Faradaic efficiency over the duration of bulk electrolysis are similar to that observed for cubic NiSe2. To understand this, Inventors comprehensively considered both the production and decomposition of the H2O2 and also computationally compare the expected catalytic properties of 2D pentagonal PdSe2 against those of pyrite phase NiSe2 with a cubic symmetry and identical (100) and (001) surfaces (
Notably, the PdSe2/GDE electrosynthesis performance is dependent on the loading and morphology of the deposited catalyst film—a PdSe2 film with higher loading on the GDE limited the Faradaic efficiency and accumulation of H2O2 during electrolysis. This was most likely due to inhibited O2 transport to the triple phase boundary. Despite this, all GDE flow cell electrosyntheses resulted in significantly more H2O2 accumulation in a much shorter time than a traditional H-cell electrosynthesis, which highlights the importance of cell design in evaluating the performance of 2e− ORR catalysts. The high H2O2 production rate (7933 mmol gcatalyst−1 h−1) at 0 V vs. RHE is notably among the highest to reach practically useful concentrations of H2O2 in a conventional flow cell device.
Post-Electrolysis and Operando Stability of PdSe2. In order to confirm the predicted stability of PdSe2 experimentally, several post-electrolysis experiments were conducted. First, in order to evaluate PdSe2's viability for direct applications where only small amounts of Pd and Se leaching into solution can be tolerated (i.e. wastewater treatment, semiconductor cleaning), we measured the amount of Pd and Se leached during electrolysis at 0.2 V vs. RHE and 0.3 V vs. RHE using inductively coupled plasma optical emission spectroscopy (ICP-OES) and found that Pd leaching was near-zero at the more active operating potential of 0.2 V vs. RHE, while Se leaching increased at the more active 2e− ORR potential (
After confirming ex-situ that the PdSe2 maintained its bulk structure and surface electronic environment before and after electrolysis, we sought to probe the dynamic electronic and local structure environments during the true operating conditions of PdSe2 during 2e− ORR using operando XAS. Interestingly, the Se K-edge XANES peak shape and edge position were unchanged from EOC up to potentials of 0 V vs. RHE (
Surface Pourbaix Diagrams. To understand the electrochemical stability of PdSe2, we calculated the surface Pourbaix diagrams of the (001) and (100) surfaces under the assumption that the surfaces are in equilibrium with bulk water. The proton-coupled electron transfer (PCET) reaction of a general adsorbate can be written as:
OmHn*+(2m−n)(H++e−)→*+mH2O(l) (1)
The electrochemical potentials of a proton and an electron pair were calculated by using the computational hydrogen electrode (CHE) method (Equation 2).
+
=0.5μH
where and
are electrochemical potentials of a proton and an electron, and μH
Clean surface is predicted for the (001) surface over a wide range of potential, which is unsurprising considering its small surface energy, indicating that this surface is much more stable than the (100) surface. Likewise, (100) surface is predicted to be bare at the applied potential although the potential window where the clean surface is predicted is narrower. Note that bond breaking is not required to generate the (001) surface because the interaction between layers is through the van der Waals interaction. Pd atoms on the (001) surface have the same coordinate number as that of bulk Pd atoms. Although the coordination number of Se atoms does not change, Se atoms of bulk PdSe2 will interact with the atoms in a neighboring layer. This will make Pd atoms on the (001) surface less active. On the other hand, Pd—Se bonds need to be broken to generate the (100) surface, which makes this surface more reactive. Because the implicit solvation model, such as VASPsol, cannot describe the competitive water adsorption, the interaction of water with PdSe2 should be taken into account, considering water adsorption energy on the (100) surface is strong and the preferential binding sites for H2O and OOH are the same. To investigate if all the Pd atoms are covered by specifically adsorbed water or not, ab-initio molecular dynamics (AIMD) was performed over 10 picoseconds by putting 30 water molecules in the 2×1 supercell of the PdSe2 (100) slab and we found that the specifically adsorbed water will occupy half of Pd atoms of the (100) surface and this can affect the adsorption energies of adsorbates, such as O*, OH*, OOH* and H*. After equilibration steps, 2 Pd atoms out of 4 Pd atoms are always covered, and desorption did not occur during AIMD. This AIMD result is consistent with the calculations along with VASPsol, which shows that the adsorption free energy of 0.5 ML (monolayer) gives −0.54 eV while that of 1.0 ML is −0.58 eV, i.e., the differential adsorption free energy from 0.5 ML to 1.0 ML is small enough (˜0.04 eV) that the thermal fluctuation can overcome. Therefore, for the (100) surface, we put 1 specifically adsorbed water molecule on Pd for all configurations. Note that both give the clean surface at the applied potential although the presence of a specifically adsorbed water molecule makes the surface more stable.
Free energy diagrams. We further calculated the free energy diagrams of 2e− ORR to investigate the activity and selectivity of the PdSe2 catalyst. Likewise, two surfaces, (001) and (100), were considered, and the CHE method was utilized to treat the electrochemical potential of a proton/electron pair. Note that there is one specifically adsorbed H2O, which corresponds to 0.5 ML, for the (100) surface. At the calculated equilibrium potential of 2e− ORR, the first PCET step is uphill by 0.45 eV and downhill by 0.34 eV on (001) and (100) surfaces, respectively, which is unsurprising considering that the (001) surface is more stable than the (100) surface, which will lead to making the OOH* state unstable. Note that the corresponding coverage used for the free energy is 0.5 ML and the stability of the OOH* state also depends on its coverage (Table 7).
Considering that the standard reduction potential of 4e− ORR (1.23 V) is higher than that of 2e− ORR (0.69 V), kinetics should be taken into account to explain the selectivity. Because the adsorbates are likely to result in 4e− ORR once the O—O bond is broken, we considered three thermal bond dissociation processes:
O2(g)→2O* (3)
OOH*→O*+OH* (4)
H2O2*→2OH* (5)
Although all these steps are thermodynamically favorable, all the activation barriers are much higher than that of the O—OH bond breaking on Pd metal (0.06 eV), suggesting that the high activation barriers can stem from the spatial separation of active sites, which will result in the high selectivity for 2e− ORR. Note that the activation free energy of the bond dissociation of hydrogen peroxide on the (100) surface is relatively small (0.24 eV), but this activation energy can be high considering the competing step is desorption from the surface, which is assumed to be barrierless.
Kinetics of PCET. In addition to thermal bond dissociation steps, we considered relevant PCET processes to investigate the selectivity for 2e− ORR:
O2(g)+(H++e−)→OOH* (6)
OOH*+(H++e−)→H2O2* (7)
OOH*+(H++e−)→O*+H2O(l) (8)
H2O2*+(H++e−)→OH*+H2O(l) (9)
To obtain activation energy for an electrochemical process, a solvated proton should be treated explicitly unlike the CHE method. Among several models for a solvated proton, such as a Zundel ion (H5O2+), the Eigen ion (H9O4+) and the water bilayer model, we used the Eigen ion as a proton donor because the use of the Eigen ion was validated for HER on Pt. Furthermore, because the purpose of this calculation lies in the comparison of activation energy, we believe it is safe to rely on the static minimum energy path and ignore the entropic effects and solvent fluctuations. In addition, we believe the trend remains the same even if the proton donor changes from a solvated proton to water. To obtain constant-potential activation barriers, the number of electrons varied so that the work function, which is related to electrode potential, is fixed. The details about this methodology can be found in J. Phys. Chem. C 2019, 123, 4116-4124, incorporated by reference, for any purpose, herein.
By using the nudged elastic band (NEB) method and the dimer method, we calculated the “raw” activation energy. The raw activation energy means the work function varies across the reaction. First, the most important step to determine the selectivity for 2e− ORR is the second PCET step (Equations 7 or 8). On the (100) surface, the step of reducing OOH to H2O2 is a barrierless step (the raw activation barrier is zero), while the step of reducing OOH to H2O(l)+O* gives a barrier of 0.44 eV. By varying the number of electrons to set the electrode potential to be −0.21 V vs SHE, which corresponds to 0.2 V vs RHE at pH 7 and constant, the potential constant activation barrier becomes 0.37 eV, i.e., the potential-constant correction term is small. This is because the change in work function is small enough and the work function is already close to the target potential. Thus, it is sufficient to consider just the “raw” activation energy to compare competing reaction steps. Similarly, others have shown that the constant potential correction term is sufficiently small, they used the average value of the potential at the initial and final states. This is particularly important for the (001) surface because there are some issues, such as a non-zero band gap and dependence of the potential of zero charge (PZC) on the interlayer distance, that make it tricky to convert raw activation barriers into potential-constant activation barriers.
Similarly, on the (001) surface, the step of reducing OOH to H2O2 becomes barrierless, while that of reducing OOH to H2O(l)+O* has a barrier of 0.23 eV. Therefore, the high selectivity for 2e− ORR can be explained by kinetics.
Additionally, we calculated the activation barrier for other steps (Equations 6 and 9). Because the further reduction of H2O2 shows a larger activation barrier than that of the first PCET step, this can explain the selectivity of H2O2 production.
pH dependence of the activity. The pH dependence of the activity of electrochemical processes like ORR has been explained by the stabilization of the OOH* state, which arises from the interaction with the electric field formed at the interface when the applied potential deviates from the PZC. Note that the OOH* state is calculated at PZC in the CHE method and the potential dependence is added as a posteriori correction via the RHE scale, where the applied potential is coupled with pH. Therefore, to decouple the electrode potential and pH, we calculated the binding energy of OOH* including the interaction of OOH with the electric field formed at the interface because the electric field formed at the interface is related to the shift in potential from the PZC (Equation 1). Others calculated the dipole moment and polarizability of each adsorbate by fitting a second-order polynomial (Equation 2) to calculations across the range of external electric fields and concluded that OOH* becomes stable as pH decreases.
In the same manner, we calculated the dipole moment and polarizability of OOH* by applying the external electric field from −0.6 V/Å to 0.6 V/Å. (
Regarding the (100) surface, the presence of specifically adsorbed water and the adsorption decrease as pH increases, i.e., it is more likely that O2 can replace the adsorbed water at high pH.
Ex-situ and In-situ Studies on the Stability of PdSe2. To confirm the predicted stability of PdSe2, Inventors conducted several post-electrolysis characterization experiments. First, in order to evaluate the viability for direct applications of the produced H2O2 where the amounts of metal and Se leaching must be low (i.e. wastewater treatment, semiconductor cleaning), Inventors used inductively coupled plasma optical emission spectroscopy (ICP-OES) to measure the amount of Pd and Se leached into the electrolytes during GDE electrosynthesis conducted with applied constant currents of −30 mA, and −60 mA, and −85 mA (
Inventors then investigated the chemical stability of the PdSe2 catalyst. Very little change was observed in the post-electrolysis Raman spectrum (
To further investigate the changes to the electronic structure of Se, Inventors conducted ex-situ X-ray absorption spectroscopy (XANES) at the Se K-edge. The similar ratio between the main white line peak and the post-white line peak (at ˜12660 eV and ˜12770 eV) confirms that Se remains anionic in character and precludes the formation of bulk Se oxides. Namely, the peak near 12770 eV was still clearly present, unlike in Se(0) powder, and cationic Se compounds (such as SeO2) which show a clearly blue shifted white line peak. Additionally, the Se K-edge extended X-ray absorption fine structure (EXAFS) spectroscopy of post-electrosynthesis PdSe2/GDE showed little change in the first two major peaks (
To confirm these stability trends on a longer timescale, we fabricated an electrode with a high loading of Nafion to prevent flooding (
Computational Methods. The Vienna Ab initio Simulation Package (VASP) interfaced with the Atomic Simulation Environment (ASE) was used for energies and geometries for all adsorbates. The core electrons are described by projector augmented wave (PAW) pseudopotentials and the Perdue-Burke-Ernzerhof (PBE) functional was used to treat exchange and correlation. Because PBE-TS (Tkatchenko-Scheffler) method gives reasonable lattice constants of bulk PdSe2, the Tkatchenko-Scheffler method was utilized to treat dispersion. Solvation effects are described by using the implicit solvation model, VASPsol. For numerical stability, the effective surface tension is set to zero, which is validated by the fact that the difference in adsorption energy between the default value (0.025 meV/Å2) and 0 meV/Å2 is less than 0.01 eV.
The Brillouin zone was sampled using 10×10×10, 10×10×1, and 10×8×1 Γ-centered Monkhorst-Pack mesh for bulk, (001), and (100) facets calculations, respectively. When larger supercells were used, the corresponding Brillouin zones were sampled.
The (001) surface of PdSe2 was modeled as a 3-layer 1×1 unit cell slab with two bottom layers fixed, which corresponds to 6 Pd atoms and 12 Se atoms. The (100) surface was modeled as 5-layer 1×1 unit cell slab with two bottom layers fixed, which corresponds to 20 Pd atoms and 40 Se atoms. For the electrochemical activation barrier calculations, 4-layer (100) surface was used for efficiency.
Each electronic self-consistent field (SCF) calculation was converged below 10−5 eV and the surface adsorbates were allowed to relax until forces became below 0.01 eV/Å. Transition states for thermal bond dissociation were searched using the nudged elastic band (NEB) method and the dimer method. Additionally, all transition states were confirmed first-order saddle points with one imaginary frequency. The free energy of H2O(l) was calculated using the experimental free energy difference between H2O(l) and H2O(g). The free energy of O2 was determined by setting it to give the experimental standard reduction potential (1.229 V) of the reaction, 0.5 O2 (g)+H2→H2O(l). The free energies of adsorbates were calculated using G=EDFT+Uthermal+ZPE−TS, where EDFT is the energy obtained from DFT calculation, Uthermal, ZPE and S are the contribution from thermal internal energy, the zero-point energy, and entropy, which are calculated under the harmonic approximation. The free energies of gas-phase/aqueous species, such as H2, H2O, and H2O2 are calculated by using Gaussian 16 with the SMD continuum solvation model. The calculated standard reduction potential of 2e− ORR is 0.83 V, which is slightly higher than the experimental value (0.69 V) and close to other computed values.
For potential-constant electrochemical barrier calculations, we included the correction term QVref, which comes from using the finite height of the system, where Q is the net charge of the DFT subsystem and Vref is the negative of the electrochemical potential of the bulk electrolyte. Note that specifically adsorbed water is not considered for the calculations of transition states. Chemicals and materials. Palladium chloride (≥99.9%) was obtained from Sigma-Aldrich. Selenourea (98+%) was obtained from Acros Organics. Carbon fiber paper (CFP, Toray: 5% wet proofing) was obtained from Fuel Cell Earth and plasma treated and briefly heated at 700° C. for 5 minutes prior to use. Sigracet 28 BC gas diffusion electrodes (GDE) were used for all GDE electrosynthesis measurements.
Materials synthesis. The PdSe2 nanoplates were synthesized by a hydrothermal method. Briefly, 1.0 mmol of PdCl2 (0.18 g) and 2.0 mmol selenourea (0.25 g) were dissolved in 9 mL of nanopure water and added to a Teflon lined 20 mL stainless steel autoclave. A piece of CFP (˜1.8 cmט2.2 cm) was also added to the autoclave. The autoclave was sealed and heated at 220° C. for 12 hours, and then cooled naturally in air. The resulting powder was washed with nanopure water and ethanol before being dried under vacuum at 60° C., and the CFP piece was similarly rinsed with nanopure water and ethanol and then dried in air. Nanoparticles of Pd4Se and Pd17Se15 were synthesized by similar methods with appropriate stoichiometric amounts of PdCl2 and selenourea.
Materials characterization. Powder X-ray diffraction (PXRD) patterns were collected on a Bruker D8 ADVANCE powder X-ray diffractometer using Cu Kα radiation with a 0.6 mm slit. Scanning electron microscopy (SEM) images were collected on a Zeiss SUPRA 55VP field emission scanning electron microscope at an accelerating voltage of 1-3 kV for imaging, and electron dispersive spectroscopy (EDS) spectra were collected on the same microscope using an accelerating voltage of 22 kV and a Thermo Scientific UltraDry EDS Detector. X-ray photoelectron spectroscopy (XPS) was performed on a Thermo Scientific K-Alpha XPS system with an Al Kα X-ray source. Raman spectra were collected on a Thermo Fisher Scientific DXR3xi Raman Imaging Microscope using a 532 nm laser with a 10 mW laser power. X-ray absorption spectroscopy (XAS) was collected at beamlines 10-ID (Pd K-edge data) and beamline 9-BM (Se K-edge data, Pd L-edge data) of the Advanced Photon Source at Argonne National laboratories in fluorescence mode with an ion chamber detector for the Pd K-edge, a Vortex silicon drift detector for the Pd L-edge, and a Passivated Implanted Planar Silicon (PIPS) detector for the Se K-edge. Pd L-edge spectra were collected in a He-purged sample chamber.
Rotating ring disk Electrode Preparation. Catalysts were drop-casted onto a rotating ring-disk electrode (RRDE, Gaoss Union, 4 mm disk diameter) which was polished successively with 1, 0.3, and 0.05 μm alumina suspensions (Allied High Tech Products) on polishing pads (Buehler, MicroCloth), and subsequently rinsed with nanopure water before brief (˜20 seconds) sonication in ethanol then quickly blown dry with N2. Catalyst dispersions typically used a 1:9 v:v mixture of Nafion solution and water with an effective catalyst loading of ˜12 μg/μL.
Rotating ring-disk electrode measurement. RRDE measurements were conducted in an undivided cell using a Bio-Logic VMP3 potentiostat. A graphite rod was used as the counter electrode and a Hg/Hg2SO4 reference electrode (CH Instruments Inc., CHI151) that was calibrated against a saturated calomel electrode (CH Instruments Inc., CHI150). The cell contained approximately 45 mL of electrolyte, which was pre-purged with Ar or O2 gas before measurement, and then the corresponding gas was kept in the headspace of the solution for the duration of the measurement. During RRDE measurements, the ring was typically held at 1.3 V vs. RHE (pre iR correction) where H2O2 reduction is diffusion limited. The H2O2 selectivity was then calculated according to the follow equation:
where N is the collection efficiency of the ring, which was calculated using a ferri-/ferrocyanide redox couple.
Bulk electrolysis measurements. Bulk electrolysis experiments were conducted in a glass H-cell with an Hg/Hg2SO4 reference electrode. The working electrode consists of the as-synthesized PdSe2 catalyst directly grown onto carbon fiber paper (CFP) cut into a 1×2 cm area, where the area exposed to electrolyte was ˜1 cm2 with the electrode contacted via a Teflon clip with a platinum contact. The other chamber, separated by a Nafion-117 membrane, contained a graphite rod which served as the counter electrode. The working chamber was pre-purged with O2 and then kept under a continuous flow of O2 gas during the measurement. Aliquots were continuously taken during the measurement and added to solutions of ˜0.4 mM Ce(SO4)2, which were then measured shortly thereafter by UV-Vis on a JASCO V-570 UV/Vis/NIR spectrophotometer, which were converted to the corresponding concentration of H2O2 via the following equation:
where F corresponds to Faraday's constant (in C/mol e−), [H2O2] is the H2O2 concentration (in mM) calculated from the previous equation, Vtot corresponds the total solution volume (in mL), and Qtot is the total charge passed through the measurement (in C).
Operando X-ray absorption spectroscopy measurements. Operando XAS measurements performed at the Advanced Photon Source were collected using a custom-built fluorescence detection H-cell where the working electrode was mounted on the front using an acrylic or aluminum face plate, with the catalyst side facing inward to the electrolyte and the bare side facing back towards the incident X-rays and the fluorescence detector. To contact the working electrode under the faceplate, a small piece of copper foil was used. A leakless Ag/AgCl reference electrode was used (and secured in the front chamber with the working electrode) and a graphite rod was used as the counter electrode in the back chamber, which was separated from the front chamber by a Nafion-117 membrane to avoid H2O2 crossover and decomposition. 2.5 mL of electrolyte was added to the front chamber for each measurement, and was kept under a continuous stream of O2 gas throughout the duration of the measurement via a port in the top of the cell. XAS spectra were continuously collected for the duration of the measurement, with each displayed spectrum representing the average of at least three raw spectra.
This application claims priority to U.S. Provisional Application No. 63/430,933 filed on Dec. 7, 2022, the content of which is incorporated by reference in its entirety.
This invention was made with government support under 1955074 awarded by the National Science Foundation. The government has certain rights in the invention.
| Number | Date | Country | |
|---|---|---|---|
| 63430933 | Dec 2022 | US |
