Effect of Urine Compounds on the Electrochemical Oxidation of Urea

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Effect of urine compounds on the electrochemical oxidation of urea using a nickel cobaltite catalyst: An electroanalytical and spectroscopic investigation Andrew F. Schranck, Randal Marks, Elon Yates, and Kyle Doudrick Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.8b01743 • Publication Date (Web): 14 Jun 2018 Downloaded from http://pubs.acs.org on June 15, 2018

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Environmental Science & Technology

Effect of urine compounds on the electrochemical oxidation of urea using a nickel cobaltite catalyst: An electroanalytical and spectroscopic investigation Authors: Andrew Schrancka, Randal Marksa, Elon Yatesb, Kyle Doudricka* Affiliations: a Department of Civil and Environmental Engineering and Earth Sciences, University of Notre Dame, Notre Dame, Indiana, USA, 46556 b Department of Civil and Environmental Engineering, Florida A&M, Tallahassee, FL, USA, 32310 *

Corresponding Author Address: Kyle Doudrick, Department of Civil and Environmental Engineering and Earth Sciences, University of Notre Dame, 156 Fitzpatrick Hall, Notre Dame, Indiana, USA, 46556 Phone: 5746310305 e-mail: [email protected] JOURNAL: Environmental Science & Technology Date: June 12, 2018

Keywords: Urine, Urea, Nickel, Electrolysis, ATR-FTIR, Wastewater

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Abstract Cyclic voltammetry (CV) and in-situ attenuated total reflectance-Fourier transform infrared (ATR-FTIR) spectroscopy were used to investigate the effect of major urine compounds on the electrooxidation activity of urea using a nickel cobaltite (NiCo2O4) catalyst. As a substrate, carbon paper exhibited better benchmark potential and current values compared to stainless steel and fluorine-doped tin oxide glass, which was attributed to its greater active surface area per electrode geometric area. CV analysis of a synthetic urine showed that phosphate, creatinine, and gelatin (i.e., proteins) had the most negative effect on the electrooxidation activity of urea with decreases in peak current up to 80% compared to a ureaonly solution. Further investigation on the binding mechanisms of the deleterious compounds using in-situ ATR-FTIR spectroscopy revealed that urea and phosphate weakly bind to NiCo2O4 through hydrogen bonding or long-range forces, whereas creatinine interacts strongly, forming deactivating inner-sphere complexes. Phosphate is presumed to disrupt the interaction between urea and NiCo2O4 by serving as a hydrogen bond acceptor in place of catalyst sites. Weak binding of urea supports the hypothesis that it is oxidized through an indirect electron transfer. Outcomes of this study contribute to the development of electrolytic systems for treating source-separated urine.

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1. INTRODUCTION In municipal wastewater treatment, nitrogenous compounds require additional treatment steps that increase operational complexity and costs. Urine is a significant contributor to this issue as it makes up only 1% of municipal wastewater volume, yet accounts for approximately 80% of wastewater nitrogen.1, 2 This positions decentralized, urine source-separation as a practical solution for reducing nitrogen removal costs.3, 4 Urea (CO2(NH2)2) is the main nitrogen compound in urine, and development of sustainable technologies for its targeted removal will be key to achieving source-separation. Electrochemical systems have emerged as a practical treatment technology for urea, presenting a potentially lower cost alternative to current denitrification methods that are limited by biological process kinetics.5 Other source-separation treatment methods for recovery of N (e.g., precipitation) require increased treatment complexity and infrastructure.1 Microbialbased electrochemical systems show promise, but they involve complex bacterial cultivation, long start-up times, and stringent working conditions.6 With the advent of nanotechnology, inorganic catalysts have emerged as favorable electrode candidates for urea electrooxidation.729 Noble metals (e.g., Pt, Rh, Ru, Ir) were the focus of many early reports,9, 22, 24-28, 30-34 but fullscale application is limited by their high cost and low-abundance. Recent advances in earthabundant catalyst development have shown that nickel, which is also the active metal in the urease enzyme, is a good alternative to noble-metals.7-16, 18, 20, 26, 29, 35-37 While promising, nickel still suffers from limitations including a high overpotential requirement (~1 V)29, 35, 38, 39 and catalyst deactivation by urea oxidation by-products.40-42 For nickel electrodes in alkaline solutions, the consensus has been that urea is oxidized through an indirect electron transfer (i.e., chemical-electrochemical). In this reaction, nickel oxyhydroxide (NiOOH) oxidizes urea while being reduced to nickel hydroxide (Ni(OH)2). Ni(OH)2 is then oxidized back to the active NiOOH via an electron transfer from the electrochemical system. Computational43 and in-situ X-ray diffraction8 evidence has been presented to support this reaction mechanism, and similar behavior has been observed for other organic compounds.44-52 Electrochemical impedance spectroscopy analysis provided contradictory evidence for the reaction mechanism, suggesting the presence of both indirect and direct mechanisms.41 A direct electron transfer indicates that urea forms inner-sphere complexes with the nickel catalyst, and this pathway would be affected by other compounds competing for reactive sites. In support of this, a decrease in current density with progressing treatment time for the electrooxidation of urea using a nickel electrode was observed, and this was attributed to catalyst poisoning or deactivation from urea oxidation by-products.5, 8 Computational analysis of the by-products suggested the adsorption/desorption of cyanate (OCN-) and carbon dioxide (CO2) products were the rate-limiting steps causing catalyst poisoning.43 Less is known about how other compounds present in more realistic matrices (e.g., urine) will affect the urea electrooxidation activity. Though, presumably, select organic compounds and ions will impact reactivity as this has been observed for platinum electrocatalysts.53-58 In this study, we investigated the effect of major urine compounds on the urea electrooxidation activity for nanostructured nickel-based catalysts. The catalysts were synthesized on various electrode substrates, and the electrochemical performance (i.e., potential and current) was investigated using cyclic voltammetry for urea-only and synthetic 4


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urine solutions. The molecular-scale sorption behavior of select deleterious urine compounds on the catalyst was investigated using in-situ attenuated total reflectance Fourier transform infrared spectroscopy. The outcomes provide support for using electrolysis to remove nitrogen from urine to improve wastewater treatment processes. 2. EXPERIMENTAL 2.1. Electrode Synthesis and Material Characterization. Complete details on the hydrothermal synthesis of nickel oxide (NiO), cobalt oxide (CoxOy), and nickel cobaltite powders (NiCo2O4) and working electrodes can be found in Section S1. Electron micrographs of NiCo2O4 and Ni powders were taken using transmission electron microscopy (TEM, FEI, Titan 80-300). Samples were suspended in isopropanol by sonication and dropcast onto copper TEM grids for analysis. The crystallinity of the powders was determined using selected area electron diffraction (SAED). The morphology of NiCo2O4 and NiO nanostructures grown on electrode supports was determined using a field emission scanning electron microscope (FESEM, FEI, Magellan 400). The oxidation states of both the synthesized powders and the supported catalysts were determined using X-ray photo-electron spectroscopy (XPS, PHI, VersaProbe II). All recorded XPS spectra were normalized to the C 1s emission (284.6 eV) and fit using PHI MultiPak software. 2.2 Electrochemical Measurements. Electrochemical experiments were conducted in a three-electrode reactor (Figure S1). The reactor consisted of a 50-mL beaker (Sigma Aldrich, Z740596-12EA), working electrode, counter electrode, reference electrode, and machined Teflon™ cap with conductive rods (McMaster-Carr, 9100K29) and set screws (McMaster-Carr, 92991A103) to secure the electrodes. The hydrothermally-prepared electrodes served as the working electrodes, Pt on fluorine-doped tin oxide served as the counter electrode, and the reference electrode was either a mercury oxide electrode (Hg/HgO in 20% KOH, Koslow Scientific Company, 5088) for alkaline pH experiments or a saturated calomel electrode (SCE, VWR, 89501-040) for pH 8.7 experiments. The potential was controlled using a Biologic SP200 potentiostat (BioLogic USA), and data were recorded and analyzed using Electrochemical (EC) Lab software. Electrochemical experiments were conducted using various electrolyte conditions. 1 M KOH (Alfa Aesar, A18854) was used for alkaline experiments and 1 M sodium perchlorate (Fisher Scientific, S360) was used for pH 8.7 experiments. The electrolyte reduces solution resistance, minimizes migration effects, and eliminates low conductivity concerns. The reported concentration of urea in human urine varies widely from 11–25 g/L (0.18– 0.42 M),59-63, while 19.82 g/L (0.33 M) is the most commonly used concentration in recent urea electrooxidation studies.9, 12, 26, 35, 42, 64 To align our outcomes with previous studies, we chose to use a urea concentration of 0.33 M. Synthetic urine was prepared as described previously65 but with 19.82 g/L urea instead of 25.0 g/L. Chemicals were used as received and included urea (Amresco, 0568), disodium phosphate (Fisher Scientific, S374-500), monopotassium phosphate (ACS grade Amresco, 0781), sodium chloride (Fisher Scientific, S271-3), ammonium chloride (BDH, BDH9208), creatinine (TCI, C0398), sodium sulfite (Spectrum, 51475), gelatin, (Amresco, 9764), and Difco nutrient broth (BD, 234000). The synthetic urine compounds and corresponding concentrations are listed in Table 1.

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Cyclic voltammograms (CVs) were used to evaluate potential (E (V)) and current density (J (μA/cm2)). Scans were performed at 10 mV/s ranging from -0.1 to 1.242 V vs. the specified reference electrode. Unless otherwise specified, all potentials in this paper are versus the Hg/HgO reference electrode (i.e., normal hydrogen electrode (NHE) + 0.1 V). Ten CV cycles were used to activate the electrode surface and develop steady-state, reproducible scans. The tenth cycle of each experiment is reported herein. CVs can vary slightly from one electrode or solution to the next based on errors in measured area and solution preparation. Experiments were conducted to show the repeatability of our cyclic voltammetry technique (Figure S2). All presented data is for the working electrodes. The onset potential (Eonset) was chosen as the potential at which significant oxidation began (i.e., when the current density curve experienced a positive inflection, as indicated by a positive change in slope of the I-V curve). Because a method for determining Eonset is not well defined in the literature, and often subjective, two other benchmarks were recorded when the previous method was not sufficient for comparing CV data. Specifically, potential values were recorded at 500 µA/cm2, denoted E500, which is similar to an approach used in the electrochemical watersplitting literature.66 E500 was used to serve as an analog to Eonset. Current density values were recorded at 0.65 V, denoted J0.65, which is based on the average potential that the peak anodic current density (Jpa) for urea oxidation occurred before diffusion limitations materialized. 2.3. In-situ Vibrational Spectroscopy. The adsorption behavior of urea and select urine compounds onto NiCo2O4 under no applied potential was analyzed using attenuated total reflectance Fourier transform infrared (ATR-FTIR) spectroscopy. Spectra were collected on an infrared spectrophotometer (Bruker, Vertex V70) using a multi-reflection, horizontal ATR cell (HATR, PIKE) equipped with a germanium internal reflectance element (IRE; surface area 6.16 cm2). A liquid nitrogen cooled mercury cadmium telluride (MCT) detector was used for all measurements. The reactor cell was monitored in real-time using Bruker OPUS software with the CHROM package. To make the catalyst film on the IRE, the catalyst was first suspended in water (1 g/L) and then bath sonicated for 30 min. 1 mL of this solution was pipetted evenly onto the surface of the IRE and allowed to dry in air overnight. Prior to analysis, unbound catalyst was removed with a gentle water rinse. Batch adsorption experiments were conducted by first filling the ATR cell with 2.5 mL ultrapure water using a syringe. All subsequent ATR-FTIR experimental solutions contained 1 M potassium hydroxide (KOH, 99.9%, Sigma-Aldrich, 306568) coinciding with electrochemical experiment conditions. After 15 min to allow for equilibration, a background scan was collected (64 scans, 4 cm-1 resolution). The solution was then drained and experimental solutions (i.e., 40 mM urea, 40 mM disodium phosphate, 40 mM creatinine) were subsequently injected into the cell. A total volume of 2.5 mL remained in the cell after injection. The IR spectrum was collected continuously throughout the injection of experimental solutions (64 scans, 4 cm-1), and it was monitored until equilibrium was achieved (~15 min). All experiments conducted in the presence of the catalyst film were replicated in the absence of a film to obtain the solution-phase spectra. Aqueous spectra were collected by the same method over the bare Ge crystal, rather than the NiCo2O4coated crystal. Peak shifts for creatinine were confirmed through replicate experiments for urea and creatinine on NiCo2O4 (Figure S3). 3. RESULTS AND DISCUSSION 6


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3.1. Materials Characterization of Electrodes. The morphology of NiO (Figure 1A) and NiCo2O4 (Figure 1B-D) consisted of flower-like nanostructures on all substrates, with the presence of small and large clusters (e.g., Figure 1C), similar to that reported previously.37 TEM micrographs (Figure S4A and B) of the catalysts synthesized in the absence of a substrate confirmed the nanosheet morphology of NiO and NiCo2O4, with planar widths of a few hundred nanometers. The SAED pattern of the NiO powder (Figure S4C) matched the (111), (220), and (311) facets of the cubic crystal structure of NiO (JCPDS 04-007-8202). The SAED pattern of the NiCo2O4 powder (Figure S4D) matched the (311) and (440) planes of the cubic crystal structure of NiCo2O4 (JCPDS 00-002-1074). These SAED patterns confirmed the nanosheets are monocrystalline. The chemical composition and oxidation state of the catalysts were determined using XPS analysis (Figure S5). For NiO, a single peak at 855.2 eV (860.8 eV satellite) in the Ni 2p3/2 and at 529.6 eV in the O 1s region matched ranges reported for NiO.67-69 For NiCo2O4, two peaks were observed for both Ni and Co in their respective 2p3/2 regions. For Co, the peak at 780.6 eV and 781.4 eV is attributed to Co(III) and 781.4 Co(II), respectively.70 For Ni, the peak at 854.6 eV and 856.0 eV is attributed to Ni(II) and Ni(III), respectively.71 This mixture of 2+ and 3+ oxidation states on the surface is in agreement with AB2O4 spinel structures.72

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Figure 1. SEM micrographs of (A) NiO on carbon paper, (B) NiCo2O4 on carbon paper, (C) NiCo2O4 on stainless steel, and (D) NiCo2O4 on fluorine doped tin oxide. 3.2. The Effect of Electrode Design and pH on Urea Electrooxidation. The addition of cobalt to nickel electrodes has been shown to improve the electrochemical activity for urea electrooxidation,37, 39 but there is no information on the electrooxidation of urea using cobalt with nickel-based catalysts on carbon paper. Figures 2A and B show the CVs for CoxOy, NiO, and NiCo2O4 catalysts on carbon paper in the absence and presence of urea and a KOH supporting electrolyte. In the absence of urea (Figure 2A), each electrode exhibited a characteristic bump in current density (e.g., NiO Eonset ~ 0.5 V), which was attributed to the 7



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oxidation of Ni(II) or Co(II) species as described previously.26, 37, 41-43, 51, 73 The sharp increase in current observed at higher potentials is due to water oxidation. In the presence of urea (Figure 2B), a new peak emerges for NiO and NiCo2O4 and is attributed to urea oxidation.26, 73 Between 0.65 V and 0.70 V, the curve reaches a maximum current density that is twentyfold higher than in the system without urea. Beyond this max, the current density decreases due to diffusion limitations in this unmixed system. Cobalt did not show appreciable activity (i.e., < 50 μA/cm2) for urea oxidation under the conditions tested (Figure 2B) and thus it was excluded from further testing. Regardless of the presence of urea, the Eonset for NiO and NiCo2O4 was between approximately 0.42–0.46 V and 0.36–0.40 V, respectively. Although cobalt was not active for urea oxidation, its addition to NiO (i.e., NiCo2O4) decreased the Eonset and increased the Jpa in all scenarios tested. This has been attributed to shifting electron densities and increased conductivity.74, 75 This moved Eonset to less positive potentials and promoted greater activity, illustrating how cobalt can improve activity while not directly acting as a urea catalyst.

Figure 2. CVs for NiCo2O4, NiO, and CoxOy in 1 M KOH and the (A) absence or (B) presence of 0.33 M urea; (C) FTO, SS, and CP substrates loaded with NiCo2O4 in 1 M KOH with 0.33 M urea; and (D) NiCo2O4 on CP in 1 M KOH (pH = 13.8) and 1 M NaClO4 (pH = 8.7) and 0.33 M urea. The inset in (A) highlights the Co region between 0.45–0.65 V. The choice of catalyst substrate, which varies in surface area and composition, affects the performance of electrochemical systems.6, 38, 76 We investigated the electrochemical activity of 8


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fluorine-doped tin oxide (FTO), stainless steel (SS), and carbon paper (CP) substrates for urea oxidation. Figure 2C shows CVs for FTO, SS, and CP loaded with NiCo2O4 in the presence of urea. Eonset and Jpa were not useful descriptors for evaluating these data due to the subjective nature of Eonset and the lack of a visible Jpa for urea due to the current increase beyond 0.5 V caused by water oxidation. Thus, we used the two benchmarks described in Section 2.2, E500 and J0.65, and these are presented in Table S3. With no catalyst present, FTO and CP electrodes performed poorly with J0.65 values below 50 µA/cm2 (Figure S6). On the other hand, a J0.65 of 5,540 μA/cm2 was observed for bare SS, which was approximately twice the J0.65 without urea present (Figure S6). This suggests SS can readily oxidize urea without additional catalysts, which is presumably due to the activity of its metals content (Fe:Cr:Ni:Mo = 67.5:17:13:2.5 wt%). The addition of NiCo2O4 onto FTO and CP increased J0.65 over 600 times (Table S3). However, NiCo2O4 addition to SS only showed a 34% increase in J0.65. (Figure 2C). The increase in electrode activity for NiCo2O4 is attributed to urea-specific reactive sites, a reduction in the overpotential for urea oxidation, and an increased available surface area due to nanostructuring. The range in E500 values was only 0.061 V for substrates loaded with NiCo2O4, meaning the substrates had little effect on Eonset. The potential at which chemical reactions begin to occur (i.e., E500), ideally should not change based on surface area, but the current density or reactivity will be affected. The superior J0.65 of the CP electrode (19,120 μA/cm2) was attributed to the greater total catalyst surface area per geometric surface area. In this study, we report current as a function of the electrode geometric surface area, i.e., the product of the length and width of the submerged electrode. This assumes the substrates have flat surfaces when in reality they have a three-dimensional structure. The FTO has a relatively flat, rough surface, the SS is a mesh, and the CP is a complex fiber network (Figure S8). The fiber structure of the CP gives it a higher total catalyst surface area per geometric area, thus it has a larger number of sites compared to equally-sized FTO and SS electrodes. While mechanistically normalizing current to the total catalyst surface area would seem ideal, reporting current as a function of the geometric area is more appropriate for realistic purposes that seek to minimize cost and footprint. The pH of an electrochemical system plays a significant role in determining electrode stability and activity, and its deviation from circumneutral affects the cost of water treatment. However, alkaline conditions are desirable for the system tested because they inhibit the production of unwanted ammonia formation by urease while also promoting urea oxidation using nickel-based electrodes. Here we compare pH 8.7 and alkaline (i.e., 13.8) pH conditions to reinforce the requirement that OH- is needed to activate the Ni(II)/Ni(III) redox mechanism for urea electrooxidation26 (Figure 2D). The J0.65 for pH 8.7 conditions in the absence and presence of urea were the same (~8 μA/cm2, Figure S9), indicating urea was not oxidized. For pH 13.8 conditions, the addition of urea resulted in a substantially higher J0.65 (19,120 μA/cm2). This inactivity under pH 8.7 conditions across the range of applied potential reinforces the need for alkaline electrolyte conditions for nickel-based electrodes. The primary reason for the performance disparity between pH conditions was attributed to the need of hydroxide to form urea-active NiOOH.51, 52 A similar pH-catalytic requirement has been observed for water oxidation.77 A secondary and related explanation for the reduced activity under pH 8.7 conditions is the increased thermodynamic barrier for oxidation reactions at lower pH as described by the Nernst equation (E = ESRP – 0.059 x pH; E = pH adjusted 9



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reduction potential, ESRP = standard reduction potential).78 The addition of OH- to the wastewater stream increases cost, complexity, and potential hazards to the treatment process, but it is required for this treatment technology. However, chemical addition is a common and accepted practice in many drinking and wastewater processes (e.g., coagulants, disinfectants), and the cost of simple electrolytes is comparable. 3.3. The Effect Urine Compounds on Urea Electrooxidation. Though urea is the primary compound in urine, other urine constituents could impact the electrode surface chemistry and electrochemical performance. The current-potential behavior of the NiCo2O4 catalyst on carbon paper was investigated for a nine-compound synthetic urine (Table S1). Additionally, each urine compound was investigated separately to determine the individual effect on urea electrooxidation. The CVs of the forward potential sweeps for these experiments are shown in Figure 3A, and the J0.65 values obtained from these are summarized in Table 1. The E500 (Figure 3B) varied minimally for all samples, ranging from 0.400 V to 0.439 V. This tight range of potential variation suggested the urine compounds did not have a major impact on Eonset. Conversely, the current density, which is an indicator of the electrochemical reaction rate, varied for all samples. The J0.65 for the synthetic urine mixture (1152 μA/cm2) was approximately three times lower than urea only samples (3145 μA/cm2), indicating one or more of the compounds had a negative impact on the urea oxidation rate. Ammonium chloride had the highest J0.65 (5154 μA/cm2), which was almost twice that of urea. This phenomenon was attributed to ammonium and not chloride, as a similar effect was not observed for sodium chloride. Under alkaline conditions, the ammonium ion will be deprotonated (pKa = 9.24) as aqueous ammonia. Previous CV analyses of ammonia have consistently shown a peak at ~0.65 V (vs. RHE),79 which overlaps with urea and thus explains why the J0.65 is greater than urea only. All compounds besides ammonia negatively impacted urea electrooxidation to varying degrees. Difco broth, NaCl, and Na2SO3 had a minor impact on the overall urea electroactivity with decreased Jpa values in the urea electrooxidation window (i.e., 0.45 to 0.70 V) of less than 30%. On the other hand, Na2HPO4, KH2PO4, creatinine, and gelatin had a more significant impact on urea electrooxidation (i.e., 0.45 to 0.70 V), with observed J0.65 values of 1316, 1318, 2176, and 512 μA/cm2, respectively, and decreased Jpa values up to ~80%. The sodium and potassium phosphate salts performed similarly to one another, with overlapping CVs between 0.45 and 0.70 V, indicating PO43-, and not Na+ or K+, was impeding urea electrooxidation through interactions with the catalyst surface and/or weak force interactions in the Stern layer. Creatinine (i.e., C4H7N3O) consists of a five-side aromatic ring with an amino and a methyl group, and these groups could potentially interact with the NiCo2O4 through weak hydrogen bonding or stronger inner-sphere complexes. Most compounds had similar current-potential behavior as urea, which includes onset potentials, peak potentials, and diffusion limitations. However, creatinine did not follow this trend; instead, a steady increase in current density with increasing applied potential was observed. This is presumably due to creatinine adsorbing to NiCo2O4 active sites (i.e., non-Faradaic current) followed by oxidation of creatinine preferentially over urea. Gelatin, which is a complex animal protein mixture, behaved similarly to creatinine, indicating strong sorption interactions with the catalyst. Due to its similar CV behavior as synthetic urine, gelatin was most likely the underlying cause of synthetic urine’s decreased activity compared to urea alone. The greater J0.65 observed for synthetic urine over 10


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gelatin was due to the additive Faradaic and non-Faradaic current obtained from active compounds (e.g., ammonia). Like gelatin, Difco broth is a heterogeneous mixture of organics (e.g., proteins), but it had only a minor impact on urea electrooxidation. This difference was attributed to the lower concentration of Difco broth (0.00016 g/L) compared to gelatin (1 g/L), which would lower its impact. To further confirm the detrimental effect of the synthetic urine components, two-hour constant voltage experiments were conducted with urea and synthetic urine at 0.65 V. Results showed a decreased current density for synthetic urine compared to the urea only experiment over the course of two hours (Figure S10). Given the likelihood of phosphate and creatinine to interfere with urea electrooxidation, further investigation was required of them to test the stated hypothesis.

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Figure 3. CV anodic sweeps of (A) synthetic urine compounds and highlighted regions detailing (B) E500 and (C) J0.65 values for select urine compounds. All solutions contained 1 M KOH. The potassium hydroxide curve represents “no urea” control, and all other solutions contained 0.33 M urea. All experiments were conducted with NiCo2O4 loaded onto carbon paper.

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Table 1. Synthetic urine compound concentrations and their J0.65 for NiCo2O4 on CP. All solutions contained 1 M KOH. All solutions contained 0.33 M urea. Concentration J0.65 Compound (g/L) (μA/cm2) Synthetic urine 1152 Urea 19.82 3145 Ammonium chloride 3 5154 Sodium sulfite 3 2926 -4 Difco broth 1.6x10 2205 Creatinine 2 2176 Sodium chloride 9 1950 Monopotassium phosphate 2.5 1318 Disodium phosphate 2.5 1316 Gelatin 1 512 3.4. In-Situ ATR-FTIR Analysis of Synthetic Urine Compound Adsorption onto NiCo2O4. In-situ ATR-FTIR was used to investigate the individual adsorption behavior of urea, phosphate, and creatinine on NiCo2O4, and the competitive adsorption behavior of urea and phosphate or creatinine on NiCo2O4. The FTIR spectra of urea in the absence of NiCo2O4 (i.e., solution-phase) agreed with previously reported major bands for urea80, 81 (Figure 4A; Table 2). Deconvolution identified three peaks between ~1700 to 1550 cm-1. The peak at 1662 cm-1 is mostly due to the in-plane symmetric NH2 bending with contributions from CO and CN stretching (i.e., δs(NH2) + ν(CO) + νs(CN)). The peak at 1632 cm-1 is due to out-of-plane asymmetric NH2 bending with contribution from symmetric CN stretching (i.e., δas(NH2) + νs(CN)). The peak at 1604 cm-1 is due to CO stretching with contributions from symmetric NH2 stretches (i.e., ν(CO) + δs(NH2) + ρ(NH2)). The peak at 1466 cm-1 is due to asymmetric CN stretching with contribution from asymmetric NH2 bending and CO bending (i.e., νas(CN) + δas(NH2) + δ(CO)). The peak at 1160 cm-1 is due to in-plane, asymmetric NH2 bending (i.e., ρ(NH2)). Urea has a resonant structure with double bond character shifting between CO and CN, and bonding of the N or O to a catalyst will shift resonance to more single bond character for CN and CO vibrations, respectively.24, 25, 34 This change in bond character upon adsorption to the catalyst would result in red- or blue- shifts for CN and CO vibrations. In the presence of NiCo2O4 (Figure 4B), only minor shifts (-3 to +3 cm-1) were observed for all peaks compared to the solution-phase peak positions. The lack of new peak development, peak splitting, or major shifts indicates that the interaction between urea and NiCo2O4 is presumably due to hydrogen bonding or long-range forces (i.e., outer-sphere complexes), and not complexation between urea N or O and the catalyst (i.e., inner-sphere complexes). However, changes in the relative absorption intensities between peaks were observed, mostly with respect to the CO stretching and NH2 bending modes. For example, the ratio between and δas(NH2) and ν(CO) increased from 0.59 to 0.85 upon interaction with NiCo2O4. This increased NH2 character is presumably due to orientation of the amide groups toward the catalyst through hydrogen bonding. This supports the hypothesized mechanism that NiOOH oxidizes urea through an indirect electrochemical-chemical pathway.41, 42, 73

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Phosphate can form inner-sphere complexes with metal (hydr)oxide surfaces,82, 83 thus it was hypothesized that the decreased current we observed in the presence of phosphate was due to site blocking. However, ATR-FTIR analysis showed that phosphate did not interact strongly with NiCo2O4. At pH 13.8, the phosphate ion was fully deprotonated (PO43-), and its solutionphase spectra matched its spectra in the presence of NiCo2O4 (Figure S11), indicating weak if any interaction occurred. The only notable difference in the spectra was a minor loss of intensity for the 1007 cm-1 peak. This can be attributed to the repulsion of the negative PO43from the NiCo2O4 surface, which is also negatively charged at alkaline pH (e.g., pHzpc = 3.5).84 Similar adsorption behavior was previously observed for TiO2.85 These results suggest there is no meaningful interaction between phosphate and NiCo2O4. Though phosphate did not form strong complexes with NiCo2O4, it did affect the adsorption behavior of urea (Figure 4C). In the presence of phosphate, only minor peak shifts in the urea peaks were observed. However, significant variations in the peak intensities (e.g., 1467 and 1165 cm-1) occurred along with slight distortion of the peak at 1662 cm-1. These anomalies, all of which involve the urea amide groups, suggest there was a change in the adsorption behavior of these functional groups when phosphate was introduced. A few possible interactions can be hypothesized. Urea can act as a hydrogen donor for hydrogen bonding between oxygens on the phosphate molecule (hydrogen acceptor) and the urea amide groups.86, 87 In the presence of phosphate and catalyst, the ratio of δas(NH2) and ν(CO) increased from 0.59 to1.35, which suggests further orientation of the amide groups toward the catalyst through hydrogen bonding between urea and phosphate. This could lessen the interaction between urea and NiCo2O4 or otherwise hinder electrochemical oxidation on the electrode surface. Alternatively, phosphate could be electrostatically blocking urea from interacting with NiCo2O4 in the Stern layer.

13



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391 392 393 394 395 396 397 398

Figure 4. In-situ ATR-FTIR spectra of 40 mM urea (A) solution-phase, (B) in the presence of a NiCo2O4 particulate film, and (C) with NiCo2O4 and 40 mM of phosphate. All experiments were conducted with 1 M KOH to mimic electrochemical pH conditions. Spectra are offset for clarity. Table 2. Urea peak locations (in cm-1) observed in this study, and a comparison to those observed or calculated previously. The bold assignments represent the major vibration. Solution

On

On NiCo2O4 and

Reported

Phase

NiCo2O4

with Phosphate

Values80, 81

δs(NH2) + ν(CO) + νs(CN)

1662

1660

1664

1663–1668

δas(NH2) + νs(CN)

1632

1629

1631

1629–1636

Peak Assignment81

399 400 401 402 403 404 405 406 407 408 409 410 411

ν(CO) + δs(NH2) + ρ(NH2)

1604

1601

1599

1597–1602

νas(CN) + δas(NH2) + δ(CO)

1466

1464

1467

1463–1473

ρ(NH2)

1160

1160

1165

1155–1160

Contrasting to urea and phosphate, creatinine, which had a complex solution-phase spectrum (Figure 5A), showed strong and complex adsorption behavior with NiCo2O4 (Figure 5B). In the presence of urea, there were numerous overlapping peaks that occurred (Figure 5C; Table S5). To date, the peak-structure relationships of the aqueous creatinine spectra have not been fully evaluated. Hydrogen bonding in crystalline creatinine creates an extremely rich spectrum that does not match the solvated creatinine species found in aqueous conditions.90 Further, relevant ATR-FTIR studies have explored the spectra only at biologically relevant pH88, 89 and not alkaline conditions (i.e., pH = 13.8) used in our study. At pH 13.8, aqueous creatinine will predominantly exist in deprotonated form (pKas = ~3.8, 12.7), the structure of which has not yet been fully determined.90, 91 For our observed spectra, peak deconvolution was not reliable due to the complexity and the numerous possible fitting outcomes. Therefore, the peaks observed in this study have not been positively identified, and a full evaluation is 14


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430 431 432 433 434 435 436 437 438 439 440 441 442


beyond the scope of the study. However, we did observe significant and obvious shifts in major peak locations and intensity along with the formation of new peaks when creatinine was analyzed in the presence of NiCo2O4 (Figure 5B). The large, complex peak cluster between 1700–1500 cm-1 is difficult to interpret, but new peaks at 1421 cm-1 and 1309 cm-1 were unique to creatinine in the presence of NiCo2O4, indicating the formation of strong, inner-sphere complexes. Upon adsorption to NiCo2O4 the solution-phase peaks at 1485 cm-1, 1462 cm-1, and 1421 cm-1 all disappear, while the peak at 1632 cm-1 is shifted to lower wavelengths (Figure 5A). These vibrations are associated with either CN or NH2, suggesting the binding occurs at the amino group and/or the heterocyclic N. Therefore, when creatinine and urea are both present, creatinine will interact preferentially with the catalyst due to its strong bonding, inhibiting or blocking urea due to its weaker interaction with the catalyst. This behavior is consistent with our electrochemical observations, which showed the presence of creatinine decreased the current generated in the urea oxidation potential window. Unlike phosphate, whose removal is not probable due to unfavorable oxidation thermodynamics, creatinine will presumably be oxidized at higher potentials, and thus removed simultaneously with urea. Increased current density at potentials greater than 0.65 V for creatinine-urea solutions compared to urea-only solutions (i.e., Figure 3A) supports this hypothesis.

Figure 5. In-situ ATR-FTIR spectra of 40 mM creatinine (A) solution-phase, (B) on NiCo2O4, and (C) on NiCo2O4 and with 40 mM urea. All experiments conducted in 1 M KOH to mimic electrochemical pH conditions. Spectra are offset for clarity. See Table S5 for peak identification. 3.5. Implications for Source-Separated Urine Treatment. These electrochemical and insitu vibrational spectroscopy analyses using a synthetic urine have revealed new information about the interaction between nickel-based electrodes and urine compounds. Phosphate, creatinine, and gelatin were major inhibitors of urea electrooxidation, hypothesized to disrupt urea oxidation through different mechanisms in the Stern Layer (Figure 6). This insight into the solid-liquid interface will be useful for furthering catalyst development. While sorption knowledge presented herein is useful, future studies should investigate the impact of applied 15



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460 461 462 463 464

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potential as changes in electron density at the catalyst surface will shift sorption behavior, as has been observed for platinum and rhodium electrodes.22, 24, 34 The high activity of carbon paper as a substrate makes it conducive for development of flow cell technologies that use membrane electrode assemblies.15, 16, 38, 92 Solution pH after urea electrooxidation remains a challenge for nickel-based electrode systems. Effluents sent to municipal treatment would not drastically impact the pH of these systems due to the low volume of urine compared to total wastewater volume. However, effluents should be neutralized to prevent transportation hazards and infrastructure damage. Depending on the goal of an electrochemical system, non-urea compounds may or may not need to be treated. The initial application of these electrolysis systems will be for N removal (and H2 generation) to reduce N burden on municipal wastewater treatment plants. Similarly, recovering direct energy from urine using a urea fuel cells is a possible early application.26, 38 In either case, downstream treatment would be required for other compounds, though a single system to achieve complete treatment is a future goal. With continued development and research focused on improving current generation while minimizing negative effects from urine compounds, electrolytic cells represent an attractive option for sustainably treating sourceseparated urine.

Figure 6. Conceptual schematic of the solid-solution interface, including interactions between NiCo2O4, urea, phosphate, and creatinine. The 0, 1, and 2 indicate the Stern plane numbers.

16


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ACKNOWLEDGEMENTS This work was made possible through financial support from Dr. Doudrick’s startup funds provided by the University of Notre Dame, and Andrew Schranck’s scholarship funding from the Patrick and Jana Eilers Graduate Student Fellowship for Energy Related Research (Center for Sustainable Energy at Notre Dame (ND Energy)), the Rothblatt Memorial Scholarship (Lake Michigan States Section of the Air and Waste Management Association (A&WMA)), and the Post Graduate Scholarship (National Collegiate Athletic Association (NCAA)). The authors thank Dr. Ian Lightcap of the Materials Characterization Facility (MCF) for aiding with the PHI VersaProbe II system. The MCF is funded by the Sustainable Energy Initiative (SEI), which is part of ND Energy. The authors thank Drs. Tatyana Orlova and Sergei Rouvimov of the Notre Dame Integrated Imaging Facility for assistance with SEM, TEM, and SAED analysis. The authors thank Rob Roberts of BioLogic for technical support with electrochemical experiments, Brent Bray and Leon Hluchota for help machinging electrochemical apparatus parts, and undergraduate students David Clark and Juan Velazquez for experimental assistance.

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Effect of Urine Compounds on the Electrochemical Oxidation of Urea

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