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Enhanced Electrocatalytic Activities by Substitutional Tuning of Nickel-based Ruddlesden-Popper Catalysts for the Oxidation of Urea and Small Alcohols Robin Forslund, Caleb Alexander, Artem M. Abakumov, Keith P. Johnston, and Keith J. Stevenson ACS Catal., Just Accepted Manuscript • DOI: 10.1021/acscatal.8b04103 • Publication Date (Web): 11 Feb 2019 Downloaded from http://pubs.acs.org on February 11, 2019

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Enhanced Electrocatalytic Activities by Substitutional Tuning of Nickel-based Ruddlesden-Popper Catalysts for the Oxidation of Urea and Small Alcohols Robin P. Forslund1, Caleb T. Alexander2, Artem M. Abakumov3, Keith P. Johnston2 & Keith J. Stevenson3* 1

Department of Chemistry, The University of Texas at Austin, 1 University Station,

Austin, Texas 78712, USA. 2

McKetta Department of Chemical Engineering, The University of Texas at Austin, 200

E Dean Keeton St., Austin, Texas 78712, USA. 3

Center for Electrochemical Energy Storage, Skolkovo Institute of Science and

Technology, 3 Nobel Street, Moscow 143026, Russia. * Corresponding author: [email protected]

Abstract/TOC Graphic

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Abstract The electrooxidation of urea continues to attract considerable interest as an alternative to the oxygen evolution reaction (OER) as the anodic reaction in the electrochemical generation of hydrogen due to the lower potential required to drive the reaction and the abundance of urea available in waste streams. Herein we investigate the effect of Sr substitution in a series of La2-xSrxNiO4+δ Ruddlesden-Popper catalysts on the electrooxidations of urea, methanol, and ethanol and conclude that activities toward the urea oxidation reaction increase with increasing Ni oxidation state. The 75% Sr―1 substituted La0.5Sr1.5NiO4+δ catalyst exhibits a mass activity of 588 mA mgox and 7.85 A


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―2

mg ―1 cmox for the electrooxidation of urea in 1 M KOH containing 0.33 M urea, demonstrating the potential applications of Ni-based Ruddlesden-Popper materials for direct urea fuel cells and low-cost hydrogen production. Additionally, we find the same correlations between Ni oxidation state and activities for the electrooxidations of methanol and ethanol, as well as identify processes that result in catalyst deactivation for all three oxidations. This demonstration of how systematically increasing Ni – O bond covalency by raising the formal oxidation state of Ni above +3 serves to increase catalyst activity for these reactions will act as a governing principle for the rational design of catalysts for the electrooxidation of urea and other small molecules going forward.

Key words: electrochemistry, Ruddlesden-Popper, perovskite, urea, electrooxidation, methanol, ethanol


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Introduction As the global demand for energy increases greater efficiency in devices used to generate and store clean energy, such as water electrolyzers for low cost hydrogen, is required to continue the pragmatic development of these key technologies. Given that the efficiencies of these devices are limited in large part by the sluggish kinetics of the oxygen evolution reaction (OER, 4OH– → O2 + 2H2O + 4e–), increased efforts have been made to reduce the overpotential for the OER in alkaline conditions with innovative catalysts.1 Recently, many have turned to the electrooxidations of small molecules as an alternative to the OER as the anodic reaction in the production of hydrogen2 or to produce electricity directly from methanol, ethanol, and urea fuel cells due to the lower overpotential required to drive these reactions.3–5 For example, while the theoretical cell potential for hydrogen generation via the OER is 1.23 V, when urea is oxidized at the anode at a theoretical potential of -0.46 V vs. SHE at pH 14 while hydrogen is evolved at the cathode at a theoretical potential of -0.83 V this window shrinks to 0.37 V, allowing for much cheaper hydrogen production and energy storage.6 Furthermore, urea is relatively non-toxic, nonflammable, and can be transported as a solid that readily dissolves in water, qualities that make it an attractive option as a means for storing chemical energy.2,7 By utilizing urea and other nitrogenous products present in the waste streams of such processes as ammonia and fertilizer production we can generate a chemical fuel while simultaneously remediating possible sources of environmental pollution.2 While the electrooxidations of methanol and ethanol have been well-studied, recent advances have been made towards catalyzing the urea oxidation reaction (UOR) using nickelbased catalysts as alternatives to precious metals and partially substituting nickel compounds with other metals has been demonstrated to reduce the onset potential for urea oxidation and


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improve catalyst activity and longevity.6,8,9 Not long ago NiMoO4 was demonstrated to be an active catalyst for the UOR and served as a model for how oxygen vacancies within the catalyst lattice can lead to an improvement in catalytic activity toward urea oxidation.10,11 Elsewhere, efforts have been made to take advantage of more sophisticated supports and possible catalystsupport interactions to boost electrocatalytic activity, such as combining Ni-based nanoparticles with carbon nanotube12,13 or tungsten carbide supports.14,15 Still others have sought to increase activities by employing extensively nanostructured materials to increase the surface area available for reaction.16–19 While many approaches have been taken to improve catalyst performance, no fundamental descriptor has been developed to aid in the systematic design of materials to increase activities towards the UOR. Botte and others have used DFT modeling to propose various steps through which the UOR may take place,20 however, the elementary mechanistic steps have not yet been experimentally elucidated. Botte and coworkers have proposed a simple, generalized mechanism for the reaction on nickel-based catalysts in which the reaction follows an indirect, electrochemical-chemical (EC’) mechanism that was developed using in-situ Raman spectroscopy.21,22 In this 6-electron pathway nickel hydroxide is oxidized to the catalytically active oxidation state of Ni3+ in nickel oxyhydroxide. Urea then reacts to form the products CO2, N2, and H2O while regenerating the nickel hydroxide catalyst in a chemical step at the anode, and hydrogen is produced at the cathode, as shown below.

Anode: 6Ni2+(OH)2(s) + 6OH- → 6Ni3+OOH(s) + 6H2O + 6e- (1) 6Ni3+OOH(s) + CO(NH2)2(aq) + 6OH- → 6Ni2+(OH)2(s) + N2(g) + 5H2O + CO2(g) (2)


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Cathode: 6H2O + 6e- → 3H2(g) + 6OH- (3) Overall: CO(NH2)2(aq) + H2O → N2(g) + 3H2(g) + CO2(g) (4)

By inspection of the above reaction it is evident that Ni in the +3 oxidation state is responsible for urea oxidation, reacting with urea and getting reduced in the process. Therefore, utilizing a material with an inherent oxidation state of Ni3+ should lead to a more efficient electrocatalytic cycle. Furthermore, the electrooxidations of other small molecules including methanol and ethanol have been shown to proceed through the same type of mechanism and should benefit from the same oxidation state augmentation.3,4 Perovskite oxides have the nominal formula ABO3 in which the A-site is most often an alkaline earth or rare earth metal and the B-site is a transition metal. After examining the above EC’ mechanism our group demonstrated how the use of the perovskite LaNiO3 in which La3+ forces Ni into the +3 oxidation state via charge compensation led to the highest activities for the UOR at that time.23 Additionally, previous work by our group also utilized perovskite materials as highly active catalysts for the OER24–26 and for Co-based perovskites27 and Ni-based Ruddlesden-Popper28 phases we demonstrated how increased M – O bond covalency upon substitution of La3+ for Sr2+ led to an increase in OER activity. Based on these studies we predicted that increasing the transition metal oxidation state past Ni3+ may lead to elevated activities for the UOR. However, upon any significant amount of Sr2+ substitution the LaNiO3 perovskite structure becomes unstable leading to phase segregation and decreased catalytic activity.29


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To avoid phase segregation and incorporate Sr2+ into a phase-pure catalyst material we turn to the Ruddlesden-Popper (RP) crystal structure. A derivative of the perovskite structure, the RP structure has the nominal formula An+1BnO3n+1 or equivalently (AO)(ABO3)n where n layers of perovskite are separated by single (AO)(OA) rocksalt slabs and as n→∞ the RP structure becomes the perovskite structure. The RP structure is able to accommodate the various elemental compositions of the perovskite structure as well as many the perovskite structure cannot.30 Herein we report the synthesis of a series of La2-xSrxNiO4+δ catalysts with the n = 1 RuddlesdenPopper crystal structure that we use to demonstrate a new governing descriptor of how systematically increasing Ni oxidation state leads to greater activities toward the UOR. Furthermore, we apply this descriptor to the oxidations of methanol and ethanol, two other reactions of interest for energy storage. While efforts to modulate the properties of metal oxide catalysts, including perovskites, to improve catalytic activity have been investigated before for the electrooxidations of methanol and ethanol,31,32 this is the first time the RP structure has been used to systematically elevate the oxidation state of Ni above +3 to improve catalyst activity for these reactions. Furthermore, the most active catalyst of the series for all three reactions, La0.5Sr1.5NiO4+δ, is the most active catalyst ever reported for the UOR to our knowledge and is very active for methanol and ethanol oxidations as well.4

Methods Catalyst Synthesis La2-xSrxNiO4+δ (LSN, x = 1 – 1.5) precursor particles were synthesized using a modified Pechini method. A and B-site nitrate salts were mixed in stoichiometric ratios


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and dissolved in water to make a solution with a total metal salt concentration of 0.1 M. Citric acid and ethylenediaminetetraacetic acid (EDTA) were added to the solution each at a ratio of 1:1 with the total metal nitrate salts. Tetramethylammonium hydroxide (TMAOH) was added to the solution until the pH reached 7.5 to ensure the EDTA lost three protons and had completely dissolved. Diethylene glycol (DEG) was then added to the solution at a concentration of 0.067 M and the solution was heated to 85°C while stirring. When heated, a dehydration reaction between the polyhydroxyl alcohol and the carboxylic acid groups of the chelates formed a polyester gel. After all of the water had been evaporated the gel was fired at 350°C to form mixed metal oxide precursor particles. This step was performed on a hot plate and not in a sealed furnace to avoid possible explosions from rapid evolution of gasses upon combustion. Finally, precursor particles were crystallized at 950°C for 5 hours under pure O2 flowing at 200 mL min-1, in a tube furnace. The furnace was first heated to 120°C in 20 minutes and held at 120°C for one hour to ensure complete evaporation of water from the tube. The temperature was then ramped to 950°C at a rate of 5°C min-1. After calcination, the furnace cooled naturally to 400°C where it was held for another 2 hours to ensure a


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maximum oxygen content had been achieved. Finally, the furnace was allowed to cool to room temperature on its own. LaNiO3 and NiO were made the same way with the exception that a calcination temperature of 700°C was used as LaNiO3 breaks down to its constituent oxides at 950°C. Catalysts were recovered and immediately stored under Ar gas to prevent catalyst surface amorphization. Electrochemical Characterization The catalyst and Vulcan carbon (VC) were each ball-milled for three minutes before being mixed together in a ball mill for three minutes in a 30:70 perovskite:carbon weight ratio. Catalyst inks were prepared by adding 2 mL of a NaOH neutralized 0.05 wt% Nafion solution in ethanol to 2 mg of catalyst powder (1 mg mL-1) and bath sonicated for at least one hour. A volume of ink (10 μL) was drop cast onto a clean 5 mm (0.196 cm2, Pine Instruments) glassy carbon electrode and dried under ambient conditions overnight. The glassy carbon electrodes were cleaned prior to drop casting by sonication in a 1:1 by volume DI water:ethanol solution. The electrode was then polished using 0.05 μm alumina powder, rinsed with DI water, sonicated in a fresh DI water:ethanol solution, and rinsed with DI water again before being dried in ambient air.


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All electrochemical tests were performed on electrodes prepared by this method, obtaining a composite catalyst loading of 51 μgtotal cm-2geo, yielding 15.3 μgoxide cm-2geo for catalysis and intercalation tests (30 wt% on carbon). Electrochemical testing was performed on a Metrohm Autolab PGSTAT302N potentiostat equipped with high speed rotators from Pine Instruments. All testing was done at room temperature in 1 M KOH (measured pH ≈ 13.7). Positive feedback methods were used to determine electrolyte resistance (6 Ω) and all data was iR compensated after testing unless stated otherwise. Each test was performed in a standard 3-electrode cell using a CH Instruments Hg/HgO (1 M KOH) reference electrode, a fritted Au wire counter electrode, and a film of catalyst ink on glassy carbon as the working electrode. All potentials are reported versus the Hg/HgO (1 M KOH) reference electrode. Quantification of Urea and Alcohol Oxidation Activities All activity testing was performed on fresh electrodes prepared via drop-casting inks with 30 wt% catalyst on VC (15.3 μgox cm-2geo). Cyclic voltammetry scans were performed at 10 mV s-1 in Ar-saturated 1 M KOH. The current at 0.6 V vs Hg/HgO - iR


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was selected from the anodic scan for comparison between materials and electrolytes. All data reported herein is the average taken from at least three tests on fresh electrodes. Powder X-ray Diffraction (PXRD) The crystal structures of all LSN materials were probed by powder X-ray diffraction using a Rigaku MiniFlex600 Diffractometer at 298 K in ambient conditions, utilizing Cu Kα radiation (1.54 Å wavelength) operating at 40 kV and 15 mA. For all tests, Ar-sealed catalyst powder was exposed to ambient air and scanned over 20 - 90° 2θ. The PXRD patterns used for Rietveld refinement were taken with a Huber G670 Guinier diffractometer (Cu Kα1 radiation; curved Ge(111) monochromator; image plate) and the refinement was performed with the JANA2006 package.33 Surface Area Analysis Nitrogen sorption analysis was performed on a Quantachrome Instruments NOVA 2000 high-speed surface area BET analyzer at a temperature of 77 K. Prior to measurements, all samples were ball milled followed by degassing in vacuum for a minimum of 12 hours at 120° F. The speciﬁc surface area was calculated using the BET method from the nitrogen adsorption data in the relative pressure range (P/P0) of 0.05 to 0.30, with a minimum R2 of 0.995 and minimum C value of 20.


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Transmission Electron Microscopy (TEM) TEM samples were prepared by crushing the crystals in a mortar in ethanol and depositing drops of suspension onto a porous carbon grid. Electron diffraction patterns, high angle annular dark field scanning TEM (HAADF-STEM) images, and energy dispersive X-ray (EDX) spectra were obtained with an aberration-corrected Titan G3 electron microscope equipped with a Super-X EDX system, operated at 200 kV using a convergence semi-angle of 21.6 mrad. Iodometric Titrations. Iodometric titrations were performed according to the referenced procedure.34–36 3 mL of deoxygenated 2 M KI solution was added to a ﬂask containing 15 – 20 mg of perovskite under an argon atmosphere. 25 ml of 1 M HCl was added and the perovskite was allowed to dissolve. This solution was then titrated to a faint golden color with a solution of ~26 mM solution of Na2S2O3 that had been pre-standardized with standard 0.1 N KIO3. Starch indicator was then added and the solution was titrated until clear, marking the end point. X-ray Photoelectron Spectroscopy


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Chemical states were characterized by a Kratos AXIS Ultra DLD XPS in 0.1 eV steps using a dwell time of 1500 ms per step with four sweeps and a monochromatic Al X-ray source (Al α, 1.4866 eV). A charge neutralizer was used for all samples. Binding energies for all spectra were calibrated relative to the adventitious carbon peak set to 284.8 eV. CasaXPS was used for all data analysis and deconvolution.

Results and Discussion Material Characterization Calcination of mixed metal oxide precursor particles resulted in nanostructured, phasepure Ruddlesden-Popper materials for all compositions as can be seen in Figure 1. For simplicity the names of the materials are abbreviated to LSNX where the letters refer to the various elements in the material and X refers to the percentage of La in the A-site. The electron diffraction patterns of the LSN materials indexed to a body-centered tetragonal unit cell with a ~ 3.8Å, c ~ 12.4Å, while Rietveld refinement (Figure S1), and powder X-ray diffraction (PXRD, Figure 1) patterns of the pristine powders confirm a body-centered tetragonal unit cell of the I4/mmm space group that is consistent with the crystal structure of a n = 1 member of Ruddlesden-Popper (RP) homologous series. Electron diffraction patterns of the LSN phase indexed on a body-centered tetragonal unit cell with a ~ 3.8Å, c ~ 12.4Å. PXRD measurements performed on LSN25 after 5 minutes and again after 1 hour of exposure to ambient conditions indicate a small amount of segregation of Sr from the RP structure at the particle surface, and the additional peaks observed indicate the formation of SrCO3 (Figure S1). The spectrum of LSN25


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taken after 1 hour of exposure to ambient conditions is almost identical to that taken after 5 minutes, indicating that the formation of SrCO3 is self-passivating and that the vast majority of the material is stable in the RP structure. These results agree with EDX mapping of high angle annular dark field scanning transmission electron microscopy (HAADF-STEM) images that indicate an almost complete homogeneity of elemental compositions throughout the particles with small regions of Sr enrichment (Figures 1, S2). The quantification of elemental ratios overall yields a compositional formula of La0.56(3)Sr1.93(7)Ni1.00(4)O4 corresponding to a phase-pure RP structure with small regions of SrO or Sr(OH)2 that make up approximately 6% of the material, regions that form SrCO3 upon exposure to air as seen in the PXRD patterns in Figure S1. Additional SEM images of the LSN series can be found in Figure S3.


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Figure 1. Physical characterization of the LSN series. a) PXRD patterns for the LSN series indicating a phase-pure n = 1 Ruddlesden-Popper structure for all four samples. b) High angle annular dark field scanning transmission electron microscopy (HAADF-STEM) image and compositional EDX maps of LSN25 showing sintered particles and homogenous distribution of the constituent elements in the RP phase with small areas or Sr enrichment in the form of SrO or Sr(OH)2. Scale bars are 200 nm c) [100] HAADF-STEM image showing perfect stacking of perovskite (BO2) and rock salt (AO)(OA) layers that propagate to the particle surface. Scale bar is 3 nm.

Precursor particles measuring between 75 nm and 250 nm sintered at high temperatures to form larger particles, and BET measurements indicate a close range of surface areas from 2.5 m2 g-1 for La0.5Sr1.5NiO4+δ (LSN25) up to 3.7 m2 g-1 for La1.5Sr0.5NiO4+δ (LSN75) (Figure S4), thus all three samples have similar particle morphology and surface area, an important characteristic if the electrochemical properties for varying compositions are to be studied and compared on a consistent basis. Iodometric titrations were used to measure the Ni oxidation state for all four LSN samples and as can be seen in Table S1, the bulk nickel oxidation state increases with the amount of Sr substitution as does the oxygen hyperstoichiometry, consistent with previous results for La2-xSrxNiO4+δ.37–39 XPS measurements performed on of all four LSN powders indicate a surface oxidation state of Ni2.19+ for LSN25 that decreases to Ni2.16+ in LSN75, consistent with a high degree of surface hydroxylation (Figures 2, S5, S6, Table S1).25 Integration of the La, Sr, and Ni peaks shows that all LSN powders are Sr-rich at the particle surface and that the A:B site ratios


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range from 3.2 in LSN25 to 3.6 in LSN75 (Figures S5-7, Table S2) which resembles previous reports by Burriel et al. wherein angle resolved XPS and low energy ion scattering (LEIS) experiments were used to demonstrate that Ni is not present at the surface of La1.67Sr0.33NiO4+δ.40 Additionally, deconvolution of the C 1s spectra indicates the presence of 4-5% CO32- species corresponding the SrCO3 detected via PXRD and EDX mapping discussed above for the LSN materials (Figures 2, S5).


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Figure 2. X-Ray photoelectron spectroscopy of the LSN series. a) Deconvolution of Ni 3p, La 4d, Sr 3d, and C 1s spectra for LSN25 after various stages of testing. b) Atomic percentage of the La in the A-site of La2-xSrxNiO4+δ materials from integration of the La and Sr XPS spectra in a) and Figure S6. c) Ni oxidation states after various stages of testing calculated from deconvoluted spectra in a) and Figure S6.

Electrochemical Analysis To evaluate the catalytic activity of the LSN series, cyclic voltammetry (CV) experiments were performed in N2-saturated 1 M KOH containing 0.33 M or 1 M urea, methanol, or ethanol at a scan rate of 10 mV s-1. Glassy carbon electrodes (GCEs) were held stationary in order to allow for more direct comparison of the activities reported herein to others reported in the literature, many of which are tested in experimental setups that utilize stationary electrodes. Although bubble formation with repeated cycling that may cause mass-transport limitations is a concern, over a single cycle this is not an issue. The majority of experiments were performed with 1 M concentrations of urea, methanol, or ethanol, however other concentrations of urea including 0.33 M were also used when quantifying urea oxidation activities to allow for comparison to other materials reported in the literature. The catalysts were supported at 30 wt% on Vulcan carbon XC-72 (VC) because while it has been previously reported that the conductivity switches from semiconducting to metal-like conductivity when x > 1 for the La2xSrxNiO4+δ

series,41 in the absence of a support the drop-cast particles only contact each other

and the GCE surface through point contacts. By supporting the LSN series on VC we minimize


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electrical contact resistances between adjacent catalyst particles and ensure good conductivity between the GCE electrode and the catalyst. Furthermore, VC is used because it lacks surface functionalities that may interact with and significantly alter the electronic structure of the catalyst, which allows us to be confident that the differences in electrochemical performance across the materials examined are due to the compositions of the catalysts themselves. Representative CVs for the UOR on all four catalysts in the LSN series are shown in Figure 3. As the Sr content and subsequent oxidation state of the Ni within the catalyst increases the current generated from urea oxidation increases as well. The most active catalyst was LSN25, ―1 containing 25% La in the A-site, which generated 1,252 mA mgox in 1 M urea at a potential of ―2 0.6 V vs. Hg/HgO and 16.695 A mg-1cmox when normalized by BET surface area while it also ―1 ―2 generated 588 mA mgox and 7.846 A mg-1cmox in 0.33 M urea making it one of the most

active catalysts ever reported for the UOR. Surface areas measured via BET gas sorption experiments and not ECSAs are used to normalize activities due to the many significant issues involved with measuring the ECSA when working with metal oxides such as large dependencies on mass loadings, differing adsorption behavior on varying material compositions, and the wide range and uncertainty in capacitance conversion factors used, all of which can result in ECSAs that vary by orders of magnitude. Figures 3 and S8 show comparisons to the LaNiO3 perovskite material referenced above as well as NiO, both supported at 30 wt% on VC and tested in the same experimental setup as the LSN series. LSN25 and LSN37, both of which contain bulk Ni oxidation states > 3+, exhibit significantly higher activities towards the UOR than LaNiO3 and NiO despite the lower mass percentage of Ni present in both RP materials. Representative CVs for LSN25 at additional urea concentrations can be found in the supporting information (Figure S9) as well as PXRD patterns to confirm the phase purity of the LaNiO3 and NiO materials


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(Figure S10). CVs of all four catalysts were performed under the same conditions but in the absence of urea. Shown in Figure 3, the redox peaks that occur in the incipient OER region on the anodic scan due to the oxidation of Ni2+ to Ni3+ appear at the same potential of 0.44 V as the onset of the UOR and serve as evidence that the reaction follows a similar process to the widely accepted EC’ mechanism discussed above. Additional comparison of LSN25 to other materials previously reported in the literature can be found in Figure S11.


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Figure 3. Catalytic activities of the LSN series toward the electrooxidations of urea, methanol, and ethanol. CVs were performed at 10 mV s-1 in N2-saturated 1 M KOH with a) 1 M urea or b) 0.33 M urea. c) CVs performed in N2-saturated 1 M KOH at a scan rate of 10 mV s-1 in the absence of a small molecule to oxidize. d) Mass activities for the LSN series as well as LaNiO3


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perovskite and NiO both supported at 30 wt% on VC in 1 M KOH containing either 0.33 or 1 M urea. CVs to measure the catalytic activities of the LSN series toward the electrooxidations of e) 1 M methanol and f) 1 M ethanol were performed at 10 mV s-1 in N2-saturated 1 M KOH. All activities reported herein were measured at 0.6 V on the anodic scan and all measurements were performed on fresh electrodes and in triplicate.

In addition to the UOR, the LSN RP series was used to perform the methanol and ethanol oxidation reactions. Similar to the trend in activities observed for the UOR, the currents at 0.6 V for the oxidations of both alcohols at a concentration of 1 M in 1 M KOH increased with ―1 increasing Sr content (Figure 3) with activities for LSN25 in 1 M MeOH of 1,554 mA mgox and ―1 1,547 mA mgox for the electrochemical oxidation of EtOH (Table S3). Furthermore, the onset

potential for all three LSN compositions for both alcohol electrooxidation reactions of 0.44 V was the same as the onset potential for the UOR as well as the Ni2+/3+ redox reactions in the absence of a small molecule, indicating that the methanol and ethanol oxidation reactions utilize a similar EC’ mechanism as for the UOR, which is in agreement with the literature.42 The amount of charge passed due to surface redox processes in the CVs displayed in Figure 3 is approximately the same across the three LSN catalyst materials with Sr contents of 50% or greater. Together with the measured BET surface areas discussed above that indicate all four LSN materials have approximately the same surface area, this demonstrates that the increase in activity towards the UOR is due to a change in the electronic structure of the catalyst and not a morphological one, such as a dependence on the amount of an amorphous Ni(OH)2 present at the outer surface of the material. The current generated from surface Ni redox reactions is much less in LSN75 than in the higher Sr content materials, as is the UOR activity, which can be explained


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by the well-documented shift in electronic conductivity of LSN RPs from semiconducting at Sr contents below 50% to metal-like conductivities when Sr makes up the majority of the A site and corroborates the influence of the bulk electronic structure on catalyst activity discussed below.43,44 To confirm that it is the changing electronic structure of the bulk material that influences small molecule oxidation activities and investigate how the LSN RP crystal structure and resulting electronic structure at the surface may change during testing, XPS measurements of LSN25, LSN50, and LSN75 were performed on unsupported materials that had been drop cast onto GCEs that were then exposed to the 1 M KOH electrolyte containing 1 M urea, and the open circuit potential of the cell was allowed to float for 10 minutes. Upon exposure to basic electrolyte, the Ni oxidation state at the surface of the catalyst is reduced for all three LSN materials examined (Figure 2). Furthermore, the La/Sr ratio increases in all three compositions so that they become La-rich as opposed to Sr-rich (Figure 2) and the carbonate peak previously present in the C 1s spectra for all three samples disappears, indicating that the SrCO3 present in in air dissolves in alkaline conditions (Figures 2, S5). To confirm that the measured current was due to urea oxidation and not the OER, rotating ring disk electrode (RRDE) experiments were performed under the same conditions as described above except with half the geometric catalyst loading (25.1 μgtotal cm-2geo) and with the electrode rotating at 1600 rpm in order to reduce bubble formation at the disk that can result in poor sensitivity of the ring electrode toward oxygen production. During these experiments the potential at the disk electrode, containing the catalyst, is swept while the Pt ring potential is held at -0.4 V vs. Hg/HgO, a potential negative-enough to reduce any oxygen that may contact it after it is evolved at the disk but not so reducing as to perform CO2 reduction. As can be seen in


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Figure S12, the ring current begins to increase at 0.69 V due to reduction of oxygen, signaling that the OER has begun to take place at the disk electrode and confirming that the current we measure at 0.6 V vs Hg/HgO is not due to oxygen evolution. Additionally, while a small amount of oxidation of the VC support may be taking place as carbon is not thermodynamically stable at such anodic potentials, were oxidation of the carbon support generating a significant amount of current then we would observe much greater activities at 0.6 V in Figure 3 where cycling was performed in the absence of urea. To investigate the stability of the catalysts both repeated cyclic voltammetry and galvanostatic constant-current tests were performed. These tests were performed under the same conditions described above except that electrodes were rotated at 1600 rpm to prevent the accumulation of bubbles on the surface of the electrode that may lead to mass transport limitations and a decrease in measured currents. Cyclic voltammetry stability tests for the UOR were performed by cycling LSN25 and LSN75 supported at 30 wt% on VC 20 consecutive times over one of two potential windows, either from 0.41 V to 0.7 V or 0.2 V to 0.7 V. The lower end of the potential window of 0.41 V was chosen by inspection of the CVs in Figure 3 where it can be seen that the onset potential for the reduction of Ni3+ to Ni2+ on the cathodic scan occurs below 0.41 V vs Hg/HgO. By cycling in this shortened window the surface of the catalyst should not be reduced due to the applied potential, but rather only due to reaction with urea or an alcohol, versus the full potential window of 0.2 V to 0.7 V that does include the Ni3+/2+ reduction peak. As can be seen in Figure 4, when cycled over the shortened window in which Ni is not reduced the activities at 0.6 V for LSN25 were stable or actually increased slightly over 20 cycles for the UOR. When the potential window is widened to include the Ni reduction peak the activities decrease with each subsequent cycle, thus when the catalyst surface is reduced by an


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applied potential it is deactivated. LSN75 displayed the same trends in activities as LSN25 over the two potential windows (Figure S13), verifying that this deactivation process occurs independently of any stability issues that may arise from Sr segregation, and these results agree with our previous work that showed how repeated cycling over the full potential window resulted in catalyst deactivation due to possible carbonate formation on LaNiO3.23 LSN25 was tested in the same manner for both methanol and ethanol oxidation and the same trends were observed (Figure S14) indicating that it is not a specific reaction product of urea oxidation that deactivates the catalyst in this manner. By keeping the potential above 0.41 V, LSN25 may remain stable and act as a highly active catalyst for the UOR as well as the electrooxidations of methanol and ethanol.


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Figure 4. Cycling stability tests for LSN25. 20 cycles were performed on LSN25 in Ar-saturated 1 M KOH with 1 M urea at a scan rate of 10 mV s-1 over a potential window of a) 0.2 V to 0.7 V. or b) 0.41 V to 0.7 V. c) The mass activities measured at 0.6 V of the anodic scan for LSN25 over 20 cycles for the electrooxidations of 1 M urea, methanol, and ethanol. d) Galvanostatic stability test for the entire LSN series in Ar-saturated 1 M KOH with 1 M urea as well as for LSN25 in 1 M KOH containing either 1 M methanol or ethanol. In these experiments the current is held for hour-long intervals and stepped upwards from 10 A g-1 up to 500 A g-1 after each hour.


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Cycling stability tests were repeated for the UOR under the same conditions as above with the exception that unsupported catalysts were used and the total mass loading was increased 10x (0.51 mg cm-2) to allow for the recovery of catalyst materials for use in post-cycling XPS analysis. As can be seen in Figure S15, concomitant increases in catalytic activity and Sr content are observed as described above for all three LSN compositions. Additionally, the same deactivation process during cycling over the full potential window was observed for all three LSN compositions tested while stable catalytic activities were observed during cycling over the shortened potential window, regardless of Sr content. The mass activities of the unsupported materials were much lower than for the materials supported at 30 wt% on VC, but this is expected due to decreased electrical conductivity of the drop cast material as well as poor catalyst utilization due to the high mass loadings that were employed in order to allow for recovery of the catalyst. As shown in Figures 2 and S5, post-cycling XPS analysis reveals a small reduction in the Ni oxidation state for all three LSN compositions compared to their uncycled counterparts that were only exposed to 1 M KOH. This reduction occurs during cycling through both the full and shortened potential windows, however, the trend of increasing surface Ni oxidation states with increasing Sr contents across the LSN series was preserved. For LSN25 and LSN50 the elemental ratios of La/Sr increase during cycling over both potential windows (Figure 2) while the total A:B-site ratios for both decrease (Figure S8), indicating some Sr leaching from the catalyst surface. On the other hand, LSN25 exhibited a small decrease in the La/Sr ratio (Figure 2) while the total A-site/B-site ratios after cycling over the two different potential windows both increased slightly, indicating that while some restructuring may have taken place there was limited Sr segregation (Figure 2). It is important to note that while the total A-site/B-site


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elemental ratios in LSN50 and LSN75 are approximately the same at all stages of testing, which show that the amount of Ni present at the surface of the catalyst is approximately equal, the catalytic activity of LSN50 for the UOR is much higher than LSN75. Overall, these results show that all LSN compositions undergo some catalyst restructuring during cycling regardless of whether the shortened window was used. While postcycling XPS analysis was unable to ascertain the cause of the catalyst deactivation observed for all three LSN compositions during cycling over the full voltage window, the oxidation state of Ni in all three compositions remained above 2+, which is the oxidation state we would expect to see if all of the catalyst material within the ~7 nm penetration depth of X-rays during XPS had restructured to form Ni(OH)2 species. Thus we can conclude that surface restructuring is limited to a thin region at the surface of the RP catalysts and that the bulk electronic structure of the catalyst still influences the electronic structure at the surface. Past work has examined how the electronic structure of LSN RP materials changes as a result of Sr substitution into the A-site and it has been shown experimentally via X-Ray characterization techniques as well as via DFT modeling that the Ni – O bonding in LSN RP materials is highly covalent. As the Sr content in LSN RPs is increased the 3d bands of the highly oxidized Ni shift downward and overlap with the O 2p band resulting in hybridization of these bands as well as pinning of the Fermi level at the top of the O 2p band.27,45 This hybridization that takes place at Sr contents equal to or greater than 50% allows for facile charge transfer through Ni – O – Ni bridges and accounts for the observed change in electronic conductivity from semiconducting to metallic.46 Additionally, due to pinning of the Fermi level at the top of the O 2p band, significant Sr substitution results in the formation of ligand holes and reduced oxygen content within the lattice, evidenced experimentally herein (Table S1) as well as


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elsewhere.40,43,44 A recent report that included DFT modeling of LSN RPs showed that the oxygen vacancy formation energy in LSN25 is negative, making the process of generating oxygen vacancies energetically favorable, and that the formation of oxygen vacancies within the perovskite layers becomes even more energetically favorable as the Sr content is further increased in Sr2NiO4.28 When LSN RPs are utilized in the electrooxidations of small molecules we believe the increased Ni – O bond covalency of highly oxidized Ni and the subsequent increasing number of ligand holes upon continued addition of Sr work to withdraw electron density from the hydroxylated, reduced Ni species near the surface of the particle, facilitating easier oxidation of the surface Ni2+ to Ni3+ and accelerating the rate at which the electrochemical step of the generalized EC’ mechanism takes place. This proposed mechanism is evidenced not only by the observed increase in catalytic activities but also by the continued increase in Ni oxidation state as the Sr content is systematically increased across the LSN series. Some amorphization of the RP structure may occur due to segregation of Sr from the outermost layer of the catalysts, however, if the amount of Ni(OH)2 at the surface was governing catalytic activity then we should not see such drastic differences in activities across the lower Sr content materials which displayed similar A-site/B-site elemental ratios, and NiO which converts to Ni(OH)2 upon exposure to alkaline electrolyte would have been the most active catalyst towards the electrooxidation of urea due to the high Ni content, not one of the least active catalysts. To further confirm the catalyst maintains its high performance over an extended period constant current tests were performed where currents were held for consecutive hour-long periods with the current increasing each hour. The currents used were stepped at intervals from 10 A g-1 up to 500 A g-1 to demonstrate the stability of the catalysts under a variety of potential


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device operating conditions. As can be seen in Figure 4, all four compositions in the LSN series were tested for the UOR using 1 M urea in 1 M KOH and LSN25 was examined in 1 M KOH in the absence of urea as well. Even in the absence of urea LSN25 was able to generate 10 A g-1 at potentials in the incipient OER region due to surface Ni2+/3+ redox and oxygen intercalation into the bulk LSN25 crystal structure, as has been discussed elsewhere.28 When 1 M urea was present all four LSN compositions were able to generate 10 A g-1 for an hour each while maintaining potentials below the onset of the OER. LSN50, LSN37, and LSN25 were able to maintain 25 A g-1 for an hour each, however, when the current was increased to 50 A g-1 in the third hour of testing LSN50 proved unable to generate enough current due to the UOR alone and the potential increased rapidly into the OER region. Only LSN25 was active enough to be able to maintain a stable potential while generating up to 500 A g-1 through urea oxidation alone. LSN25 was then subjected to the same stability test but in the presence of 1 M methanol and ethanol. As can be seen in Figure 4, LSN25 was able to maintain a stable potential while generating as much as 500 A g-1 through the electrooxidations of methanol and ethanol as well, highlighting the remarkable activity and catalytic stability of LSN25 for all three reactions. Similarly to the cycling stability experiments performed above, as long as the catalyst surface is not reduced due to an applied potential that causes deactivation, LSN25 performs as a remarkably active, stable catalyst for the electrooxidation of small molecules in alkaline conditions.

Conclusions A series of nanostructured La1-xSrxNiO4+δ (1 ≤ x ≤ 1.5) materials were synthesized via a modified Pechini synthesis followed by calcination at high temperatures. The result was a collection of phase-pure n = 1 Ruddlesden-Popper catalysts that were utilized to investigate the


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effect of increasing Ni oxidation state on the electrochemical oxidations of urea, methanol, and ethanol. Previous work by our group, inspired by looking at the EC’ mechanism described above in equations 1-4 for the UOR, in which Ni3+ can be seen as the active form of Ni showed how Ni3+ in LaNiO3 outperformed other Ni2+-based catalysts for the electrooxidation of urea. Building on that work we have demonstrated how systematic incorporation of Sr2+ into the Perovskite derivative n = 1 Ruddlesden-Popper crystal structure in place of La3+ steadily increases the bulk oxidation state of Ni past +3 leading to greater activities towards small molecule oxidations. LSN25, in which La makes up 25% of the A-site and Sr makes up 75%, displayed the highest activities for all three electrochemical oxidations, most notably generating ―1 1,252 mA mgox in 1 M urea. Furthermore, we discovered that this material can maintain its

activity over long periods of time and at very high current densities but that reduction of the catalyst surface at sufficiently cathodic potentials leads to deactivation towards these reactions. This is the first time the fundamental guiding principle of increasing Ni oxidation state past Ni3+ to increase catalyst activity towards the urea oxidation reaction has been shown and the activities of LSN25 towards the electrooxidations of urea, methanol, and ethanol are among the highest ever reported.

ASSOCIATED CONTENT Supporting Information. Additional information regarding detailed experimental procedures, additional supporting data, SEM and TEM images, catalyst surface area analysis via and nitrogen sorption curves is available free of charge via the Internet at http://pubs.acs.org. AUTHOR INFORMATION Corresponding Author


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*E-mail: [email protected]; [email protected] (K.J.S.). Author Contributions All Authors have given approval to the final version of the manuscript. R.P.F. performed all electrochemical data collection and analysis. C.T.A. contributed XPS measurements of the LSN materials. A.M.A performed TEM and ED measurements and crystallographic analysis. K.J.S. and K.P.J. assisted in constructing the manuscript. Notes The authors declare no competing financial interests ACKNOWLEDGMENT Financial support for this work was provided by the R. A. Welch Foundation (Grants F1529 and F-1319).

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ACS Catalysis

Coefficients in La0.50Sr0.50CoO3−y. J. Electroanal. Chem. Interfacial

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