Synthesis and Electrocatalytic Activity of Ammonium Nickel Phosphate

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Cite This: Inorg. Chem. XXXX, XXX, XXX−XXX

Synthesis and Electrocatalytic Activity of Ammonium Nickel Phosphate, [NH4]NiPO4·6H2O, and β‑Nickel Pyrophosphate, β‑Ni2P2O7: Catalysts for Electrocatalytic Decomposition of Urea Andrew G. Meguerdichian,† Tahereh Jafari,† Md. R. Shakil,‡ Ran Miao,‡ Laura A. Achola,‡ John Macharia,‡ Alireza Shirazi-Amin,‡ and Steven L. Suib*,†,‡,§ †

Institute of Materials Science, University of Connecticut, U-3136, 97 N. Eagleville Road, Storrs, Connecticut 06269, United States Department of Chemistry, University of Connecticut, U-3060, 55 N. Eagleville Road, Storrs, Connecticut 06269, United States § Department of Chemical & Biomolecular Engineering, University of Connecticut, U-3222, 191 Auditorium Road, Storrs, Connecticut 06269, United States ‡

S Supporting Information *

ABSTRACT: Electrocatalytic decomposition of urea for the production of hydrogen, H2, for clean energy applications, such as in fuel cells, has several potential advantages such as reducing carbon emissions in the energy sector and environmental applications to remove urea from animal and human waste facilities. The study and development of new catalyst materials containing nickel metal, the active site for urea decomposition, is a critical aspect of research in inorganic and materials chemistry. We report the synthesis and application of [NH4]NiPO4· 6H2O and β-Ni2P2O7 using in situ prepared [NH4]2HPO4. The [NH4]NiPO4·6H2O is calcined at varying temperatures and tested for electrocatalytic decomposition of urea. Our results indicate that [NH4]NiPO4·6H2O calcined at 300 °C with an amorphous crystal structure and, for the first time applied for urea electrocatalytic decomposition, had the greatest reported electroactive surface area (ESA) of 142 cm2/mg and an onset potential of 0.33 V (SCE) and was stable over a 24-h test period.

1. INTRODUCTION

position of urea to carbon dioxide, CO2; nitrogen, N2; and hydrogen, H2 are2,10−12

Excretion of nitrogen-containing organic compounds is common among biological organisms such as fish (ammonia, NH3), reptiles/snakes and birds (uric acid, C5H4N4O3), and mammals (urea, CO(NH2)2), where urea in industrial wastewater and fertilizers leads to research in the potential for electrocatalytic decomposition of urea into hydrogen, H2, gas for clean energy applications. 1−3 Additionally, urea is economical, stable, and nonhazardous to handle in large quantities for development in Direct Urea Fuel Cells (DUFC).4,5 Alternative methods of decomposing urea include urea hydrolysis, use of enzymes or bacteria, and strong oxidants or adsorption, which have disadvantages due to high cost and lack of safety at high temperatures, high pressures, and materials not selective for the overall process.1,6 Current industrial processes attempt to decompose a commercial urea solution, AdBlue (5 M, 32.5%); however, no available technology decomposes AdBlue.7 Challenges in catalyst design are the use of expensive and precious metals (Pt/C; Ni/C, anode; and Ag/C, cathode), to prevent coordination of a catalyst to urea (Pd, Pt, Cr, Fe, Zn, Cu), and to perform the reaction at standard conditions.7−9 Nickel is the primary transition metal for urea decomposition with desirably porous, highly physicochemical and electroactive surface areas. Chemical equations for electrocatalytic decom© XXXX American Chemical Society

− Anode: Ni(OH)2(s) + HO(aq) ⥂NiOOH(s) + H 2O(l) + e−

(1) − CO(NH 2)2(aq) + 6HO(aq)

→ N2(g) + 5H 2O(l) + CO2(g) + 6e− − Cathode: 6H 2O(l) + 6e− → 3H 2(g) + 6HO(aq)

(2) (3)

Overall: CO(NH 2)2(aq) + H 2O(l) → N2(g) + 3H 2(g) + CO2(g)

(4)

Besides representing half-cell reactions in the anode, cathode, and overall electrochemical analyses and in situ X-ray diffraction, studies on the Ni(OH)2/NiOOH (Ni2+ ⇄ Ni3+ redox transition, respectively) catalysts during urea decomposition have been completed to describe diffusion and kinetics in the reaction.11,13,14 Morphology and particle size15,16 were previously considered as nanoribbons with significantly higher electroactive active surface area (2114 cm2/g) than bulk nickel Received: October 23, 2017

A

DOI: 10.1021/acs.inorgchem.7b02658 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry hydroxide (97 cm2/g).17 Nanowires are also active in degradation of urea, hydrogen peroxide, and mixtures of the two substances18 along with gas chromatography analyses to quantify gaseous products N2, O2, and H2 produced.12 Composites such as (a) nickel/electrochemically reduced graphene oxide (ERGO) and nickel/graphene,10,19 (b) nickel with carbon nanofibers20 or sponge,21 and (c) nickel/tungsten carbide (WC)22 or nickel with carbon nanotubes/WC have shown activity.23 Properties such as weak adsorption on the catalytic surface and electron transfer were important for the catalytic reaction in formation of NiOOH.22,23 Increased surface area, particle morphology, and catalyst stability were critical to the carbon-nanofiber/sponge- and graphene-containing catalysts.10,19−21,23 Multiple metals have been incorporated along with nickel to enhance electrocatalytic activity in decomposing urea such as cobalt,24−29 lanthanum,30 manganese,31 molybdenum,32 and zinc.29 Heterometal oxide systems such as BiOX−TiO2 have been applied to urea electrocatalytic decomposition reactions.33 Presence of Ni3+ in LaNiO3 demonstrating the Ni2+ ⇄ Ni3+ redox couple,30 surface area, and conductivity in NiMoO4· H2O32 and nonstoichiometry in addition to average oxidation state of manganese in nickel manganese oxides31 were determined as contributing factors toward the reaction. Zinc and/or cobalt in the nickel material has been demonstrated as another effective method for increasing catalytic activity to increase surface area24,28 and electrical conductivity and facilitate electron transfer to form Ni3+ active species25,26,28 and access to nickel active sites.27,29 Inorganic phosphates include a broad class of materials such as ammonium metal phosphates, metal orthophosphates, metal pyrophosphates, and metal phosphides. Applications of inorganic phosphates are dyes and pigments,34 magnetism,35 electrocatalysis in the hydrogen evolution reaction (HER)36 and fuel cell technology,37 catalytic decomposition reactions,38 and storage of electrochemical energy.39−42 Furthermore, the synthesis and thermal properties of transition metal phosphates has led researchers to develop applications of these materials.43−47 The literature demonstrates that nickel is an important material for electrocatalytic degradation of urea into CO2, N2, and, potentially, H2 gas products. Reaction kinetics, an electrochemical mechanism, and preliminary surface properties of catalysts used for degradation of this biomolecule have been discussed and investigated.11,13,14 However, in this work, ammonium nickel phosphate hexahydrate, [NH4]NiPO4· 6H2O, is precipitated via in situ preparation of ammonium hydrogen phosphate, [NH4]2HPO4, at room temperature, and [NH4]NiPO4·6H2O is calcined at 300, 600, and 900 °C (βNi2P2O7), which are applied for their potential electrocatalytic activity in the decomposition of urea.

added dropwise to the ammonium hydrogen phosphate solution. The precipitate was vacuum-filtered with copious amounts of acetone and stored in a vacuum desiccator. Calcination of the [NH4]NiPO4·6H2O sample was completed using a tube furnace at a heating rate of 10 °C·min−1 to 300, 600, or 900 °C and held isothermally for 2 h in atmospheric air. Preparation of Electrolyte, Electrode Ink for Electrocatalytic Decomposition of Urea. Electrolyte and analyte solutions for electrocatalytic testing of the synthesized catalysts were prepared as follows: in a 100 mL volumetric flask, a 1 M potassium hydroxide (Sigma-Aldrich, reagent grade, 90%, flakes) KOH solution was prepared by dissolving 5.61 g of KOH in 100 mL of deionized water. The urea solution was prepared in a 100 mL volumetric flask by dissolving 0.1 mol of KOH and 0.033 mol (1.98 g) of urea in order to prepare a 0.33 M urea solution in 1 M KOH. Approximately 2 mg of Cabot carbon black and 2 mg of catalyst were measured in a sample vial. Distilled deionized water (800 μL) and ethanol (200 μL) were added to the sample vial and sonicated for 5 min. Then, 75 μL of Nafion was added, and the solution was sonicated for 30−60 min. The prepared ink, 10 μL, was dropped onto the electrode and allowed to dry for approximately 24−48 h. Electrochemical analysis conditions included voltages of 0−0.6 at 0.01 V·s−1 with a sampling interval of 0.001 V. Quiet time was 2 s, and sensitivity was 0.01 A·V−1.

3. CHARACTERIZATION OF [NH4]NIPO4·6H2O AND β-NI2P2O7 X-ray diffraction (XRD) measurements were obtained on a Rigaku Ultima IV instrument using Cu Kα radiation (1.5406 Å). Beam voltage is 40 kV with a beam current of 45 mA. Measurements were conducted at a 2°·min−1 measurement rate and between 5° and 75°. Thermogravimetric analysis−mass spectrometry (TGA-MS) measurements of decomposition of the prepared phosphates were made with a Netzsch TGA-MS instrument under an air/argon atmosphere (flow rate = 45/5 mL·min−1, respectively) and heated at 10 °C·min−1 to 900 °C. Scanning energy microscopy (SEM) analyses were done on a field emission instrument (FE-SEM, JEOL 6335F) along with energy dispersive X-ray spectroscopy (EDS) to study the particle morphology and presence of N, Ni, P, and O, respectively. N2 adsorption measurements, using the BET method, were conducted on a Quantachrome Autosorb IQ instrument and were degassed for 3 h at 150 °C, except for [NH4]NiPO4·6H2O, which was degassed at room temperature for 3 h. Pore size distributions and analyses were conducted using DFT methodologies using the adsorption branch of the obtained isotherm. Analyses were performed between relative pressures 0 to 1, and the adsorbent was treated as a zeolite. High-resolution transmission electron microscopy (HRTEM) analysis was conducted on an FEI Talos F200X microscope operating at 200 kV equipped with an EDS. To test each sample for urea decomposition, cyclic voltammetry (CV) measurements were conducted on a CHI Electrochemical Workstation equipped with a saturated calomel electrode (SCE) reference electrode during measurements at a urea concentration of 0.33 and 1 M KOH electrolyte. Conditions for CV analyses included scanning from 0 to 0.6 V at a 0.01 V·s−1 scan rate and sampling interval of 0.001 V. Chronoamperometry measurements were conducted on the [NH4]NiPO4·6H2O calcined at 300 °C sample (with and without 0.33 M urea) at a constant voltage of 0.5 V for 24 h to measure the change in current for decomposition of urea over this time period. Electroactive surface areas (ESA) were calculated using previously reported methods.17,23 Detection of gas desorbed from the electrocatalytic decomposition of urea was completed

2. EXPERIMENTAL SECTION Preparation of Ammonium Nickel Phosphate Hexahydrate, [NH4]NiPO4·6H2O, and β-Nickel Pyrophosphate, β-Ni2P2O7. All chemicals and reagents in this research analysis are used as received without any purification. In general, [NH4]NiPO4·6H2O was prepared34 by dissolving 4.871 g of Ni(NO3)2·6H2O (Alfa-Aesar 98%) in 37.5 mL of distilled deionized water in a 150 mL beaker. In a 500 mL round-bottom flask, 3.5 mL of ammonium hydroxide (J.T. Baker, 28.0−30.0% Baker Analyzed, ACS Reagent), 1.125 mL of phosphoric acid (J.T. Baker, Baker Analyzed, ACS Reagent), and 100 mL of distilled deionized water were dissolved. The nickel nitrate hexahydrate solution is poured into a 250 mL separatory funnel and B

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Figure 1. XRD pattern of (A) as synthesized [NH4]NiPO4·6H2O, (B) [NH4]NiPO4·6H2O calcined at 300 and 600 °C, and (C) β-Ni2P2O7. HRTEM images: (D) [NH4]NiPO4·6H2O and (E) β-Ni2P2O7. Patterns are indexed according to JCPDS Card #01-086-1866 ([NH4]NiPO4·6H2O) and JCPDS Card #01-074-1604 (β-Ni2P2O7).

4. RESULTS

via chronoamperometry at a constant voltage of 0.5 V and 0.33 M urea in 1 M KOH. The electrochemical cell utilized for the decomposition of urea consisted of a two-compartment cell system, in which both compartments were purged with argon for 10 min. Gas chromatography (GC) measurements were conducted using a SRI 8610C instrument with argon as a carrier gas and thermal conductivity detector (TCD). The pressure of the carrier gas was 20 psi and consisted of a SupelCo packed, 6-ft, 13X, molecular sieve column along with a SupelCo packed, 6-ft, 13X, Silica Gel column. The GC method for gas analysis included an initial temperature at 40 °C, which was held for 2 min; then the temperature was increased and heated at 15 °C· min−1 to 220 °C, which was held isothermally for 20 min. A sample of gas before the electrocatalytic decomposition of urea was obtained as a blank run. The electrocatalytic reaction was performed for several hours with periodic sampling from cathode and anode sample compartments of the electrochemical system.

XRD analyses (Figure 1A−C) were carried out to determine crystal structures of the nickel samples. The XRD patterns (Figure 1A and C) demonstrate that the [NH4]NiPO4·6H2O and β-Ni2P2O7 were crystalline and corresponded to the ICDD JCPDF Cards #01-086-1866, Quality: S and #01-074-1604, respectively. However, the samples prepared at intermediate temperatures (300 and 600 °C), in Figure 1B, showed a very broad peak between 2θ = 20° and 40° and were classified as having low crystallinity. HR-TEM analyses, in Figure 1D and E, demonstrated the presence of diffraction planes corresponding to the crystalline phases of [NH4]NiPO4·6H2O and β-Ni2P2O7. TGA-MS of the [NH4]NiPO4·6H2O and β-Ni2P2O7, in Figure 2A and B, demonstrated that the [NH4]NiPO4·6H2O lost approximately 50% of its mass while being heated to 900 °C. Species desorbed from the sample included water and ammonia where most of the decomposition occurred before 200 °C. However, the β-Ni2P2O7 was thermally stable up to C

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Figure 2. TGA-MS analyses of (A) [NH4]NiPO4·6H2O and (B) βNi2P2O7.

900 °C and had minimal mass loss and no detectable species being desorbed from the sample. Figure 3A−D shows SEM analyses revealed that the [NH4]NiPO4·6H2O particles at varying calcination temperatures consisted of blocks of geometric shaped particles bunched together. EDS analyses in Figure 3A−D demonstrated that nitrogen, nickel, phosphorus, and oxygen were distributed throughout the analyzed areas, when applicable. Figure 3B−D illustrate that the morphology of the particles remains the same as the calcination temperature increased from 300 to 900 °C. N2 adsorption analyses in Figure 4A and B, using the BET method, were used to demonstrate the type of porous structure and shape/contour of the pores in the synthesized materials. Figure 4A illustrates that the isotherms for the prepared [NH4]NiPO4·6H2O at room temperature to 600 °C were reflective of a type II isotherm with a H3 hysteresis loop. However, with increasing calcination treatment to 900 °C, the isotherm transformed to a type III with a H3 hysteresis loop.48 The [NH4]NiPO4·6H2O sample demonstrated a type II isotherm in which an inflection point in the curve at P/P0 of 0.1 was observed where a monolayer of nitrogen was adsorbed on the material surface.48 The hysteresis type can also be characterized as a type H3 in which the adsorption of the material is based on a nonrigid type of particles. As the calcination temperature increased, the isotherm of the [NH4]NiPO4·6H2O sample gradually changed into a type III isotherm.49 After calcination to produce the β-Ni2P2O7, a nonporous sample was obtained that is reflective of a type III isotherm in which a monolayer of nitrogen is not pronounced. The pore size distribution in Figure 4B, for the as-prepared [NH4]NiPO4·6H2O demonstrated the sample to have a broad range of pore sizes. While the shape of the nitrogen isotherm provides information about the porosity and adsorption/desorption nature of the sample, physicochemical parameters of surface area, pore size, and pore volume can be deduced from the analyses and are summarized in Table 1. The [NH4]NiPO4· 6H2O sample had a surface area of 56 m2/g, pore size of 2.6 nm, and pore volume of 0.10 cc/g. As the calcination temperature increased, the sample became nonporous with surface areas decreasing to 23 m2/g (300 °C), 26 m2/g (600

Figure 3. SEM images and EDS elemental mapping of [NH4]NiPO4· 6H2O (A) without calcination treatment and calcined at (B) 300 °C, (C) 600 °C, and (D) 900 °C (β-Ni2P2O7).

°C), and 7 m2/g (900 °C). The DFT fitting error varied between 0.426% and 1.789% (Table 1), demonstrating a good correlation between the analysis method and the obtained data. Electrochemical measurements, particularly CV, were conducted on the synthesized samples (Figure 5A−D) to demonstrate the electrocatalytic decomposition of urea. The analyses were conducted with and without urea, and a redox D

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[NH4]NiPO4·6H2O) and were incorporated into the summary of Table 1. Electrochemical data, including CVs, ESA, and onset potential values, together suggested the 300 °C sample should be further investigated for stability measurements in electrocatalytic decomposition of urea, which had the greatest ESA value (142 cm2/mg). GC analyses conducted on the gaseous products produced during urea electrocatalytic decomposition were analyzed, and N2 and O2 were observed. However, the presence of H2 was also most likely qualitatively observed as demonstrated in Figure 7A and B where bubbles of H2 gas were produced on the cathode of the electrochemical cell.

5. DISCUSSION The electrocatalytic decomposition of urea using [NH4]NiPO4· 6H2O and its amorphous (300 and 600 °C) and crystalline (βNi2P2O7) calcined materials were studied as gaseous species were produced. Two critical aspects of this electrocatalytic reaction, including kinetics and diffusion, were previously addressed in the literature.11,20,21 While the porosity of materials has been justified for this reaction, the synthesis and application of nickel nanomaterials is a continuous endeavor.50 XRD and TGA-MS analyses together demonstrated that the [NH4]NiPO4·6H2O material decreases in crystallinity (eventually becoming amorphous) as ammonia and water are being desorbed from the sample. A potential observation can be made that the ammonium and water species in the [NH4]NiPO4· 6H2O preserve the crystalline nature of the material, but with increasing temperature, the structure slowly collapses. However, the crystalline β-Ni2P2O7 was not obtained until higher temperatures were achieved (∼900 °C), where the thermally stable (evidenced by TGA-MS) phase was present. The morphology of the particles and elemental composition, as demonstrated through SEM and EDS analyses, provided information that increased heat treatment did not affect the geometric shaped particles of [NH4]NiPO4·6H2O. In fact, heat treatment at 900 °C proved that the morphology of the βNi2P2O7 remained relatively similar to that of the [NH4]NiPO4· 6H2O at room temperature. Moreover, N2 adsorption analyses provided fascinating clues into the surface porosity of the samples with increasing heat treatment. The [NH4]NiPO4· 6H2O sample exhibited an interesting phenomenon of cavitation, in which a change in the volume adsorbed at P/P0 = 0.44−0.52 in the desorption curve was noted. The phenomenon observed, cavitation, is a networking effect in which the methods of filling and emptying of pores are not equivalent. However, as ammonia and water were removed

Figure 4. N2 adsorption (A) isotherms for prepared [NH4]NiPO4· 6H2O samples at room temperature, 300, and 600 °C and β-Ni2P2O7. (B) Pore size distribution of [NH4]NiPO4·6H2O.

peak indicating the redox transition Ni2+ ⇄ Ni3+ was observed. In general, the prepared catalyst materials demonstrated the Ni2+ ⇄ Ni3+ redox peaks when urea was not present in solution. However, slight reduction peaks were observed in Figure 5A−C (as prepared [NH4]NiPO4·6H2O and samples calcined at 300 and 600 °C), whereas the sample in Figure 5D demonstrated a more obvious oxidation event occurring in the CV analysis. After conducting electrochemical measurements, ESA values were calculated for each sample for degrading urea. The asprepared [NH4]NiPO4·6H2O and β-Ni2P2O7 had ESAs of 115 cm2/mg and 44 cm2/mg, respectively. The amorphous, intermediate calcination temperature, 300 and 600 °C, [NH4]NiPO4·6H2O samples demonstrated relatively high ESA values of 142 cm2/mg and 130 cm2/mg, respectively. The onset potential, another measure of activity of a catalyst for electrocatalytic reactions, was measured by determining the potential at which 10% of the maximum current density was obtained for that sample. Overall, onset potentials varied between 0.309 V (β-Ni2P2O7) and 0.338 V (as prepared

Table 1. Summary of Electrochemical and Physicochemical Properties for [NH4]NiPO4·6H2O at Room Temperature, 300, 600, and 900 °C (β-Ni2P2O7) sample [NH4]NiPO4·6H2O [NH4]NiPO4·6H2O calcined at 300 °C [NH4]NiPO4·6H2O calcined at 600 °C β-Ni2P2O7

onset potential (V)a surface area (m2/g)b 0.338 0.33 0.316 0.309

56 23 26 7

pore size (nm)

pore volume (cc/g)

DFT fitting error

ESA (cm2/ mg)c

2.6 non-porous non-porous non-porous

0.10 N/Ad N/Ad N/Ad

0.426% 0.573% 0.539% 1.789%

115 142 130 44

a

Onset potential is defined as the potential (volts, V) at which 10% of the current density was obtained. Reference electrode used was SCE. Degassing conditions for nitrogen adsorption were 3 h at 150 °C, except for [NH4]NiPO4·6H2O, which was degassed at room temperature for 3 h. c Electroactive surface area (ESA) calculated using ESA = Q/mq where Q = charge for reduction of NiOOH to Ni(OH)2 = integral under oxidation curve of current vs time, m = mass of catalyst used in electrochemical testing = 2 mg, q = charge to form monolayer of Ni(OH)2 = 257 μC·cm−2 based on one electron transfer from Ni3+ to Ni2+.1,12,17,23 dSample is nonporous. Therefore, pore volume is not reported for this sample. b

E

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Figure 5. Electrochemical analyses (CV curves) of (A) [NH4]NiPO4·6H2O at room temperature and calcined to (B) 300 °C and (C) 600 °C. Sample D is [NH4]NiPO4·6H2O calcined to 900 °C, which produced the β-Ni2P2O7 crystal phase. Black curves represent the urea sample, whereas red curves represent no urea.

Figure 6. Chronoamperometry analysis of [NH4]NiPO4·6H2O calcined at 300 °C. Black curve represents urea sample, whereas red curve represents no urea.

from the sample, cavitation became less pronounced, and instead, a nonporous material was obtained. The activity of the samples for electrocatalytic decomposition of urea was then evaluated using cyclic voltammetry and chronoamperometry. The presence of nickel along with redox peaks indicating the Ni2+ ⇄ Ni3+ reaction demonstrated the prepared catalysts’ electroactivity. However, the amount of cavitation decreased as calcination temperature increased from [NH4]NiPO4·6H2O (room temperature) to β-Ni2P2O7 (prepared at 900 °C). The room temperature [NH4]NiPO4·6H2O and high temperature β-Ni2P2O7 demonstrated good activity toward decomposing urea. An inference could be made that the pore structure of cavitation prevented access to the Ni2+ sites in [NH4]NiPO4·6H2O, thereby making the sample not the most active. In addition, the nonporous structure of β-Ni2P2O7 also did not allow for urea to be adsorbed into the material and make maximum contact with the Ni2+, which was oxidized to Ni3+ sites in the sample. An explanation for the amorphous 300 and 600 °C samples should be proposed for their improved activity toward the decomposition of urea when compared to crystalline [NH4]NiPO4·6H2O and β-Ni2P2O7. While N2 adsorption analyses in Figure 4 demonstrated to some degree that cavitation occurs for both samples, the overall porosity was becoming more

Figure 7. (A) Overall image of two-cell compartment electrochemical apparatus and (B) more focused qualitative image of potential H2 bubble at the cathode during the electrocatalytic decomposition of urea.

nonporous; nevertheless, the 300 and 600 °C samples were able to adsorb and follow similar isotherms in adsorbing nitrogen. The XRD patterns demonstrated that these amorphous samples lacked chemical species such as NH4+, PO4−3, H2O, and P2O7−4, present in the crystalline phosphate phases. A potential explanation to correlate the activity with the isotherm analyses is that the 300 and 600 °C samples were able to expose Ni2+ sites more completely as other chemical species not participating in electrocatalysis were removed via heat treatment. As a result, when these samples were placed in a highly basic environment and voltage was increased, the Ni2+ F

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Inorganic Chemistry sites were subsequently oxidized to the highly active Ni3+ species, the active center for the decomposition reaction. Facile access and exposure of Ni2+ sites to eventually convert to Ni3+ active centers and decreased cavitation permit urea’s contact with Ni3+ and allowed for increased activity (ESA) in the amorphous samples as compared to the crystalline phases of [NH4]NiPO4·6H2O and β-Ni2P2O7. ESA values obtained for other nickel based catalysts in the literature, without using supports, have indicated that the amorphous nickel phosphate material prepared at 300 °C (142 cm2/mg) has relatively greater activity than in these studies. Electrodeposition was used in the preparation of Ni and Ni−Zn electrocatalysts and revealed ESA values of 12.1 cm2/mg and 67.9 cm2/mg, respectively.29 Variations in morphology have demonstrated that the ESA would vary as shown for nickel nanowires (ESA = 79.1 cm2/mg) and nickel hydroxide nanoribbons (ESA = 2114 cm2/g or 2.114 cm2/mg).12,17 Moreover, the GC results provide information about the gaseous species being desorbed from the sample during the electrocatalytic reaction. The expected products from the decomposition of urea include N2, CO2, and H2. While atmospheric air contains N2 gas, the electrochemical cell was purged with argon prior to conducting measurements. Therefore, the only source of nitrogen species was from urea. While CO2 would be expected in the decomposition of urea, this species was not observed due to the lack of O2 in the electrochemical cell, and instead, a saturated inert atmosphere where oxidation would be unlikely was present. Therefore, the carbonyl in urea did not undergo the expected oxidation to CO2, but rather, the species potentially formed was carbon monoxide, CO (Supporting Information, Figure S1). Literature data suggest that CO2 would not be observed in GC analyses as this chemical product becomes soluble in the electrolyte solution.12 Moreover, the solubility of CO is significantly lower than CO2 in water, which would support the idea that CO would likely be in the headspace of the electrochemical cell.51 Additionally, while H2 was not directly observed in the GC, its presence was most likely qualitatively identified at the cathode side of the electrochemical cell. Factors such as concentration and inaccessibility of H2 at the tip of the electrode in the electrochemical cell affected its potential to be observed in GC measurements. Lastly, GC analysis of the cathode gas products revealed the presence of O2 and N2, which are predicted to originate from atmospheric O2 and N2 contamination in the open lab environment. More specifically, the origin of these samples could have arisen from environmental or atmospheric N2 in the gas syringe, which was used to obtain a gas sample from the anode and cathode electrochemical cells, followed by bringing the syringe to the GC instrument in the open lab environment (O2 and N2 from atmosphere). The N2 also has another plausible explanation for its appearance in the GC analysis of the cathode compartment as a small amount could have originated from the anode cell by diffusion, where N2 was being produced. Since the electrochemical chambers between the anode and cathode cells were open and connected, the N2 could have potentially traveled to the other compartment by diffusion. The present nickel based catalyst material prepared at 300 °C and findings in the electrocatalytic decomposition of urea present a critical aspect in catalytic design of nickel nanomaterials for production of H2 for clean energy applications.

6. CONCLUSIONS The successful synthesis via direct precipitation of crystalline [NH4]NiPO4·6H2O and β-Ni2P2O7 (at 900 °C) along with amorphous materials at 300 and 600 °C were successfully applied toward electrocatalytic decomposition of urea. A thorough analysis of characterization data revealed the [NH4]NiPO4·6H2O exhibited a unique surface phenomenon of cavitation along with the calcined materials at 300 and 600 °C. The degree of cavitation eventually decreased as the βNi2P2O7 exhibited a nonporous material surface. Electrocatalytic analysis revealed that the amorphous sample prepared at 300 °C demonstrated relatively low onset potential for the decomposition of urea, compared with the others in this study. In addition, the ESA for the 300 °C sample was 142 cm2/mg, which was the largest electroactive surface area from the synthesized samples. The increased accessibility and availability of Ni2+ active sites in the 300 °C sample allowed for increased adsorption and diffusion of urea to demonstrate excellent electrocatalytic activity for decomposition of urea.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.7b02658. Gas chromatography (GC) analyses of observed gases from anode and cathode electrochemical compartments (PDF)

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AUTHOR INFORMATION

Corresponding Author

*Tel.: +1 860 486 2797. Fax: +1 860 486 2981. E-mail: steven. [email protected]. ORCID

Steven L. Suib: 0000-0003-3073-311X Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the U.S. Department of Energy, Office of Basic Energy Sciences, Division of Chemical, Biological, and Geological Sciences under Grant DE-FGO286ER13622.A000. In addition, the authors would like to acknowledge Islam Mosa and Mohamed Nabil (Dr. James Rusling’s research laboratory at UConn) and Habiba Tasnim (Dr. Steven Suib’s research laboratory at UConn) for helpful discussions about electrocatalysis and cyclic voltammetry measurements. The HR-TEM studies were performed using the facilities in the UConn/Thermo Fisher Scientific Center for Advanced Microscopy and Materials Analysis (CAMMA). Financial support for this work was provided by the Institute of Materials Science (IMS) and the use of the Bioscience Electron Microscopy Laboratory of the University of Connecticut (UConn) and use of Nova NanoSEM and Oxford EDS purchased through NSF grant #1126100.

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DOI: 10.1021/acs.inorgchem.7b02658 Inorg. Chem. XXXX, XXX, XXX−XXX