Ammonia electro-oxidation on nickel hydroxide: phases, pH and poisoning

Abstract

Nickel hydroxide is a leading alternative to platinum group metals for electrocatalysis of the ammonia oxidation reaction (AOR), an important process for energy conversion and environmental remediation. Nevertheless, the dependence of AOR electrocatalysis on the different crystalline phases at the electrode surface (α-Ni(OH)₂/γ-NiOOH vs. β-Ni(OH)₂/β-NiOOH) has never been investigated. Herein, the crystalline β-Ni(OH)₂ and the disordered α-Ni(OH)₂ were synthesized and characterized by XRD, HRSEM, and Raman and FTIR spectroscopies. The respective electrocatalytic activity of the two phases was analysed at a broad range of ammonia concentrations (0.01–2 M) and pH values (11–13). Both phases electrocatalyze the oxidation of NH₃ to N₂, as proven by online mass spectrometry, but the α-Ni(OH)₂/γ-NiOOH couple is more active. At high ammonia concentrations (>1 M), surface poisoning by adsorbed NH₃ prevents access to OH⁻, leading to less NiOOH formation, lower AOR currents, and suppression of the OER side reaction. The poisoning is strong and irreversible on α-Ni(OH)₂, as confirmed by soaking experiments. The difference in ammonia adsorption and electrocatalytic activity between the α-Ni(OH)₂ and β-Ni(OH)₂ emphasizes the importance of understanding the phase space of nickel hydroxide electrodes when designing low-cost electrocatalysts for the nitrogen cycle.

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Introduction

The electro-oxidation of ammonia is important in the fields of energy and environmental remediation.^1–5 Ammonia is a common pollutant present in wastewater, due to its extensive exploitation in industry and agriculture. The electro-oxidation of ammonia is used in electrochemical sensors^6–8 and in wastewater remediation.⁹ In the energy sector, the (alkaline) ammonia oxidation reaction (AOR, reaction (I); E° for the opposite reduction reaction = −0.77 V vs. SHE) is used in direct ammonia fuel cells,^10–12 or to produce hydrogen fuel from ammonia-containing water for energy storage applications.^5,13–16 The production of hydrogen by ammonia electrolysis is theoretically more cost-effective and less energy-consuming than water electrolysis. Furthermore, wastewater remediation can be coupled with hydrogen production.^9,11 Ammonia offers high energy density (3000 W h kg⁻¹) and higher hydrogen storage capacity (17.7 wt%) than other hydrogen carriers, without releasing CO₂.⁹ Moreover, infrastructure for ammonia synthesis and distribution already exists, due to its role as a fertilizer.

2NH₃ + 6OH⁻ → N₂ + 6H₂O + 6e⁻ (I)

However, the AOR is a complex reaction, involving the transfer of 6e⁻ and 6OH⁻. Its sluggish kinetics drive a search for better electrocatalysts. The best AOR electrocatalysts are based on platinum group metals (PGMs),^13,17–23 although they are costly and scarce. Moreover, PGMs are quickly deactivated during the AOR due to surface poisoning by adsorbed atomic nitrogen.^24–26 Thus, PGM electrocatalysts have not yet found broad commercial use.²⁷ Non-PGM alternatives include materials based on copper,^28–31 silver,³² iron^29,33 and cobalt.³⁴ Lately, the abundant and low-cost nickel oxyhydroxides, formed on nickel surfaces under alkaline and anodic conditions, demonstrated promising activity towards the AOR.^11,35 Synthetic methods such as alloying^{14,30,36–43} or nano-structuring^11,44–47 were developed to reduce the large overpotentials, which range around 1.3–1.4 V on Ni-based phases (compare to ∼0.4 V on Pt²⁶).

Nickel is a fascinating and complex electrode material, first proposed for batteries in 1899.^48,49 Despite over a century of electrochemical studies and broad applications in Ni–MH batteries,⁵⁰ Ni–Cd batteries,⁵¹ and supercapacitors,⁵² the electrochemistry of the Ni(OH)₂ surface remains elusive. The reason lies in the variety of crystalline phases of Ni(OH)₂ and NiOOH, which show different degrees of hydration and intercalation,⁵³ and interconvert through multiple routes (Fig. 1) and in different locations inside the electrode. The nickel hydroxide polymorphs range from crystalline brucite-structured β-Ni(OH)₂ to a family of badly crystalline structures denoted under the umbrella term α-Ni(OH)₂.⁵⁴ The latter phase can intercalate water and anions, be turbostratically disordered (imperfect layer orientation), and can contain hydrogen bonding and point defects.^55,56 Upon oxidation, the crystalline β-Ni(OH)₂ transforms into crystalline β-NiOOH (reaction (II)), while the α-Ni(OH)₂ becomes γ-NiOOH (reaction (III)).

β-Ni^II(OH)₂ + OH⁻ → β-Ni^IIIOOH + e⁻ + H₂O (II)

α-Ni^II(OH)₂ + OH⁻ → γ-Ni^III/IVOOH + e⁻ + H₂O (III)

Fig. 1 Bode diagram of the limiting cases of divalent and trivalent NiO_xH_y phases.^55–57

The α/γ couple has larger interlayer distances than the β/β couple, allowing ion intercalation and better ionic conductivity. Moreover, the γ-NiOOH phase can have higher average oxidation than 3+ on the Ni,⁵⁶ leading to higher electronic conductivity than β-NiOOH.⁵⁸ The α/γ couple is more catalytic towards the urea oxidation reaction (UOR),⁵⁹ while the β/β couple is considered better for the oxygen evolution reaction (OER)⁶⁰ and the hydrogen evolution reaction (HER).⁶¹ For the AOR, Łuczak and Lieder have hypothesized that the α/γ couple could be more catalytic towards the AOR.³⁵ Importantly, the α/γ couple is less stable thermodynamically, aging to the β/β couple over time, or by the effects of high temperature, alkalinity or continuous electrochemical cycling.^55,56

Several studies focused on the mechanism of the AOR, mostly for Pt and other noble metals,^{19,25,26,62–65} though only a few examined Ni(OH)₂/NiOOH surfaces. Kapalka et al. demonstrated that ammonia oxidation on Ni/Ni(OH)₂ occurs only after the surface has been oxidized to nickel oxyhydroxide,⁶⁶ as was later supported by in situ Raman studies.^67–69 The chemical oxidation of ammonia on NiOOH was further found to be spontaneous, as demonstrated by detecting nitrite (NO₂⁻) formed on NiOOH under no bias.⁷⁰ Two possible mechanisms were proposed for the AOR (in analogy for the oxidation of organic molecules): a direct electron transfer from the ammonia to the oxyhydroxide film (reactions (IV) and (V)), or an indirect oxidation where the formed oxyhydroxide is chemically reduced back to the hydroxide by an ammonia molecule (reaction (VI)).⁶⁶

Ni(OH)₂ + OH⁻ → NiOOH + H₂O + e⁻ (IV)

NiOOH + NH₃ → NiOOH(NH₃)_ads → NiOOH + products + ne⁻ (V)

NiOOH + NH₃ + mOH⁻ → Ni(OH)₂ + products (VI)

Kapalka et al. have found the Ni^III/Ni^II reduction currents were independent of ammonia concentration, suggesting that Ni(OH)₂ is not formed chemically by reaction (V), but rather a direct electron transfer occurs. However, only a narrow range of ammonia concentrations (<0.15 M) was studied.⁶⁶ Others have reported similar behaviour, where the current increases with ammonia concentration, up until 100–200 mM, then the current slightly decreases.^40,71 A mechanism switch to direct oxidation was proposed for higher NH₃ concentrations. The authors hypothesized that the mechanism switch is due to the saturation of active sites with ammonia, reducing the local OH⁻ concentration and blocking NiOOH formation.⁷¹ Thus, many questions remain open. Among them, importantly, is the effect of the nickel hydroxide phases (α/γ vs. β/β) on AOR electrocatalysis, which has never been investigated.

We now report a detailed investigation of AOR electrocatalysis and its dependence on the phase of nickel hydroxide. For the first time, the electrochemistry of α-Ni(OH)₂ and β-Ni(OH)₂ are compared directly, over a broad range of conditions relevant to energy and environmental applications, namely ammonia concentrations from 0.01 to 2 M, and pH values from 11 to 13.^13,72,73 We found that the α-Ni(OH)₂/γ-NiOOH couple is more active towards the AOR throughout the concentration range, although the β-Ni(OH)₂/β-NiOOH is also AOR-active. Both hydroxide phases are poisoned at high ammonia concentrations, leading to lower currents and delays in the onset potentials. The poisoning is stronger and irreversible in α-Ni(OH)₂, probably due to stronger adsorption energies. The effects of pH on AOR further confirm the poisoning, as do ammonia soaking experiments.

Experimental procedures

Material synthesis

α-Ni(OH)₂ and β-Ni(OH)₂ were synthesized according to Chakrabarty, Offen-Polak et al.⁵⁹ NiSO₄·6H₂O (Alfa Aesar, 98%) and NaOH (Alfa Aesar, 98%) were mixed in 35 mL of DI water in either 1 : 2 molar ratio to give β-Ni(OH)₂ (9.8 and 19.6 mmol, respectively) or a 3 : 1 molar ratio for α-Ni(OH)₂ (9.8 and 3.3 mmol, respectively). For β-Ni(OH)₂, the mixture was stirred for 30 min until it became homogeneous, transferred to a 50-mL Teflon-lined stainless-steel autoclave and placed in an oven at 120 °C for 24 hours. For α-Ni(OH)₂, the mixture was stirred at room temperature for 3 hours. The resulting green powders were collected by centrifugation, washed with DI water repeatedly (until pH reached 7–8), and air-dried at 50 °C.

Material characterization

Powder X-ray diffraction (XRD) was performed in a Rigaku SmartLab X-ray diffractometer with Cu-Kα radiation (λ = 1.5418 Å), scan rate of 1° min⁻¹. A Raman spectrometer (LabRam, 1800 grating) was used with 532 nm laser excitation. Attenuated total reflectance Fourier transform infrared (ATR-FTIR) transmittance spectra were collected on a Bruker Tensor 27 spectrometer. High-resolution scanning electron microscopy (HRSEM) was performed in secondary electrons mode on a Zeiss Ultra+ microscope, at a 4 kV acceleration voltage. The Ni concentration in the solution was measured by ICP-OES 5100 (Agilent Technologies).

Electrochemical testing

The ink solution was prepared with 1 mg of the catalyst (either α-Ni(OH)₂ or β-Ni(OH)₂) dispersed by 30 min sonication in a solution containing 300 μL deionized water and 180 μL ethanol and 20 μL of 20 μL Nafion® 5 wt% dispersion (Alfa Aesar). 10 μL of the ink solution was applied by drop-cast on a polished glassy carbon electrode (5 mm diameter) and dried at 50 °C. The catalyst loading was 0.02 mg or 0.1 mg cm⁻². Electrochemical measurements were performed by using a Biologic VSP multichannel potentiostat. Reversible hydrogen electrode (RHE) was used as reference electrode and a graphite rod as counter electrode. The electrolyte contained 0.1 M KOH and ammonia water solution (25% v/v) at the desired concentrations, along with 0.5 M K₂SO₄, required to maintain high ionic conductivity. The pH was monitored before and after the addition of ammonia, and pH was not changed by more than 0.2 pH units, except for 2 M ammonia, where it deviated by 0.4 units. For measurements at lower pH, (NH₄)₂SO₄ was used instead for ammonia source, and the pH was monitored and adjusted accordingly by addition of KOH. Before a measurement, the electrolyte was purged for 30 min by Ar at 25.0 ± 0.1 °C and flowed continuously above the solution during the experiments. For the Fe-free experiments, electrolytes were cleaned by the Trotochaud procedure⁷⁴ in a 0.1 M KOH solution. Then, the K₂SO₄ and ammonia were added to the Fe-free solution. Cyclic voltammetry (CV) was recorded at a scan rate of 10 mV s⁻¹. Before measurements, each electrode was cycled in the non-faradaic potential range to improve wetting (0.8 V to 1.2 V vs. RHE, 10 cycles, 100 mV s⁻¹). Additionally, each electrode went through 10 activation cycles before ammonia addition (0.8 V to 1.55 V vs. RHE, 10 cycles, 10 mV s⁻¹). Measurements were repeated at least 3–6 times for each concentration and pH combination, for each material, and all presented trends were shown to be consistent. Rotating ring-disk electrode (RRDE) experiments were performed at each ammonia concentration, where CVs were measured (after the initial activation) at 1600 rpm rotation, E_disk = 0.8–1.9 V vs. RHE and E_ring = 0.6 V vs. RHE. A small and constant baseline ring current was subtracted.

Differential electrochemical mass spectroscopy (DEMS). A HPR-40 (Hiden, UK) instrument was used with a custom-built cell (Fig. S1, ESI†). The DEMS probe was fitted with a microporous Teflon membrane. 20 μL of an ink solution were drop casted on a GC electrode. The ink solution was made from 1.6 mg of the catalyst (either α-Ni(OH)₂ or β-Ni(OH)₂) and 0.7 mg of carbon black (30 wt%; BP200, Beyond Battery). The ink was dispersed by 30 min sonication in 300 μL deionized water, 180 μL ethanol and 20 μL of 20 μL Nafion® 5 wt% dispersion. The carbon black acted as a conductive additive to increase the current density, boosting the N₂ production above the MS detection limit. The electrode was placed in the cell with 5 mL of the 0.1 M KOH and 0.5 M K₂SO₄ electrolyte. The probe was inserted in the solution, keeping 200–300 μm distance to the electrode surface. The electrode went through 10 cycles of activation, before adding 1 M ammonia. The DEMS measurements were conducted operando while cycling the potential from 0.8–1.55 V vs. RHE at 5 mV s⁻¹. The CVs were mathematically stacked to reduce noise.

Electrochemical surface area (ECSA). Two methods were used and then compared, due to the difficulties in determining the ECSA for nickel hydroxides.⁷⁵ First, the capacitance was estimated by the Watzele method, involving electrochemical impedance spectroscopy (EIS).⁷⁶ The electrode was scanned to from 0.8–1.4 V vs. RHE at a scan rate of 1 mV s⁻¹, while rotating at 400 rpm. EIS spectra were recorded with a frequency range of 30 kHz to 10 Hz, with an amplitude of 10 mV, at every 10 mV from 1.4 to 1.7 V vs. RHE. At least 60 data points were collected for each spectrum. The EIS spectra, taken at the initial stages of OER, were fitted to the suggested equivalent circuit (Fig. S9, ESI†). The adsorption capacitance (C_a) of the oxygenated intermediates were derived, and then correlated to the surface area. Those were compared to values derived from double layer capacitance, as measured by CV in the non-faradaic region.

Results and discussion

Phase characterization

To investigate the AOR electrocatalysis on the different phases of nickel hydroxide, both α-Ni(OH)₂ and β-Ni(OH)₂ were synthesized. The synthetic procedure and the characterization are fully described in our previous study for a broad range of crystallinities.⁵⁹ The crystallinity of the nickel hydroxide materials was determined by X-ray diffraction (XRD). Indeed, the sharp peaks of the β-Ni(OH)₂ match the diffraction pattern of this phase (JCPDS 14-0117, Fig. 2a). Moreover, the XRD confirms a crystalline and ordered β-Ni(OH)₂, while the α-phase is amorphous, with barely visible broad peaks. The morphologies of the catalysts were characterized by high resolution scanning electron microscopy (HR-SEM, Fig. 2b and c). The α-Ni(OH)₂ phase consists of irregular agglomerates (hundreds of nanometers in size), while the β-Ni(OH)₂ powder is composed of uniform hexagonal platelets (∼60 nm).

Fig. 2 Material characterization of the α-Ni(OH)₂ and β-Ni(OH)₂ catalysts, as prepared: (a) XRD, with the β-Ni(OH)₂ lines marked (JCPDS 14-0117). (b) and (c) HRSEM micrographs of the two materials. (d) Raman spectra and (e) ATR-FTIR spectra of the two materials.

The structures of the α-Ni(OH)₂ and β-Ni(OH)₂ phases were further characterized by Raman spectroscopy (Fig. 2d) and Fourier transform infrared (FTIR) spectroscopy (Fig. 2e). Raman peaks at 318, 445, 872 and 3590 cm⁻¹ are typical for β-Ni(OH)₂, and at 460 and 3656 for α-Ni(OH)₂. In the latter, a broad band at 3000–3600 cm⁻¹ arises from the high hydration of the structure.^56,77 Furthermore, strong sulphate intercalation peaks appear for α-Ni(OH)₂ at 603, 985 and 1060 cm⁻¹, indicating a disordered and open structure allowing intercalation.^78–80 The FTIR spectra confirm the phase assignments, with fingerprint peaks of β-Ni(OH)₂ at 340, 418, 506 and at 3630 cm⁻¹, and for α-Ni(OH)₂ at 374, 455 and 673 cm⁻¹.^78,81 The hydration of the α-Ni(OH)₂ phase appear as a broad band at about ∼3400 cm⁻¹. In contrast, the sharp peak at 3630 cm⁻¹ corresponds to free hydroxyl groups in β-Ni(OH)₂, as the closed-packed crystalline structure does not permit hydrogen bonding with intercalated water.⁸¹

Ammonia oxidation electrocatalysis

The AOR electrocatalytic activity of α-Ni(OH)₂ and β-Ni(OH)₂ was investigated by CV, before and after the addition of ammonia. Without ammonia (Fig. 3a), both α-Ni(OH)₂ and β-Ni(OH)₂ show oxidation and reduction peaks related to the reversible nickel hydroxide ⇌ nickel oxyhydroxide reaction (reaction (VI)). The voltametric waves on crystalline and well-defined β-Ni(OH)₂ are sharper than the CV wave on the disordered α-Ni(OH)₂ material. Furthermore, the α/γ couple has an earlier onset potential than the β/β couple (1.38 V and 1.42 V vs. RHE, respectively), and its oxidation produces a larger total charge. When ammonia is added (0.75 M, Fig. 3b), both catalysts show higher current densities for oxidation, and a different shape of the wave, suggesting that ammonia is indeed oxidized. The oxidation onset potentials are also delayed, to 1.39 V for α-Ni(OH)₂ and to 1.45 V for β-Ni(OH)₂. The electrocatalytic activity of nickel hydroxide towards the AOR is further confirmed by detection of N₂(g) by differential scanning mass spectrometry (DEMS, Fig. 3c, d and Fig. S1, ESI†). In this experiment, the DEMS probe is placed above the electrode, measuring the gaseous species emitted during cell operation. Note that the DEMS experiment only provides qualitative information on the evolution of gaseous species and could be used to determine faradaic efficiencies. The rise in N₂ production follows the voltametric ammonia oxidation peak, for both phases, taking into account a 5 second delay due to the bubble formation, release, and detection time (Fig. 3d). This proves, for the first time, that both NiOOH surface phases (γ and β) electrocatalyze the oxidation of NH₃ to N₂. The fact that the AOR starts only after the NiOOH formed confirms the hypothesized role of Ni³⁺ in the AOR electrocatalysis.⁶⁶ This is analogous to UOR and OER, where the catalytic reaction follows the transition to NiOOH.^60,82 Consequently, the use of pure nickel hydroxide as an anode in a direct ammonia fuel cell is impractical due to this high overpotential. This distinction is crucial, as was recently stressed for direct urea or alcohol fuel cells.⁸³ The activity of nickel hydroxides towards the AOR is about 5 times inferior to its UOR activity.⁵⁹ The poor performance of the pure nickel hydroxides explains the extensive focus on alloying and doping studies.^4,29,35–42 Nevertheless, the fundamental differences between the nickel hydroxide phases are crucial to explore for the future design of doped catalysts, and for general understanding of the nitrogen cycle.

Fig. 3 Cyclic voltammetry (CV) of α-Ni(OH)₂ (red) and β-Ni(OH)₂ (black) in (a) 0.1 M KOH + 0.5 M K₂SO₄ electrolyte and (b) with 0.75 M NH₃ added. (c) Differential electrochemical mass spectrometry (DEMS) of the N₂ signal produced (black, top) during CV at 5 mV s⁻¹ from 0.8 V to 1.55 V (red, bottom) with 1 M NH₃ added. (d) Enlarged portion of the DEMS experiment, focusing on the onset of N₂ evolution; the traces are averaged over 7 measurements, to improve signal-to noise ratio.

A side-reaction of O₂ evolution could contribute to the anodic wave. However, there is no OER peak in the CV up to 1.55 V, in the absence of ammonia (Fig. 3a). When ammonia is added, the catalytic wave becomes sharp and harder to deconvolute (Fig. 3b), so to rule out the OER, we performed RRDE for both α-Ni(OH)₂ and β-Ni(OH)₂ (Fig. 4). The disk-deposited Ni(OH)₂ electrodes were cycled up to a high potential (1.9 V), where OER is expected (reaction (VII)), while the ring potential was held at 0.6 V_vs. RHE, to detect the produced O₂ by reducing it (reaction (VIII)). The RRDE experiments were conducted at 1600 rpm at four ammonia concentrations of 0, 0.1, 0.75 and 1.5 M.

4OH⁻ → O₂ + 2H₂O + 4e⁻ (VII)

O₂ + 2H₂O + 4e⁻ → 4OH⁻ (VIII)

In the absence of ammonia, OER starts only above 1.55 V vs. RHE, on both α-Ni(OH)₂ and β-Ni(OH)₂. Upon increasing the ammonia concentration, the ring current diminishes, and the ring onset potential is delayed. In other words, as ammonia concentration increases, less and less O₂ can be detected. In fact, OER is barely occurring in the presence of 1.5 M ammonia, as evidenced by almost no O₂ ring current even at high overpotentials. The suppression of OER in the presence of NH₃ proves that both α-Ni(OH)₂ and β-Ni(OH)₂ are highly selective towards the AOR, and suggests a competition over the Ni³⁺ active sites between NH₃ and OH⁻. Interestingly, such suppression was observed on oxidized nickel foam (that is, covered in Ni(OH)₂/NiOOH-covered).⁸⁴

Fig. 4 RRDE experiments at 1600 rpm, with different added ammonia concentrations (0, 0.1, 0.75 and 1.5 M). The disk potential was scanned between 0.8 to 1.9 V, while the Pt ring potential was set to 0.6 V to detect O₂. The disk and ring currents are presented vs. the disk potential for (a) α-Ni(OH)₂ and (b) β-Ni(OH)₂.

AOR at different ammonia concentrations and pH values

To understand how the Ni(OH)₂/NiOOH phases behave at different ammonia concentrations, we cycled both α-Ni(OH)₂ and β-Ni(OH)₂ electrodes in 0.01 to 2 M NH₃ (Fig. 5a). This broad range covers the ammonia concentration expected in different applications, from waste-water remediation (0.01–0.07 M)⁷² to direct ammonia fuel cells and ammonia electrolyzers for H₂ production (0.5–5 M).^13,73

Fig. 5 The oxidation current density at E = 1.54 V for both α-Ni(OH)₂ (red circles) and β-Ni(OH)₂ (black squares) over a wide range of (a) ammonia concentrations, and (b) pH of the electrolyte. Each data point is an average of at least 3 measurements. A representative CV for each data point appears in Fig. S2–S5 (ESI†).

With the increase in ammonia concentration, the current rises linearly, up until 0.75 M, for both α-Ni(OH)₂ and β-Ni(OH)₂. This confirms that both phases can drive AOR electrocatalysis. Comparing the performance of the two catalysts, α-Ni(OH)₂ is clearly a more active electrocatalyst than β-Ni(OH)₂, with higher current densities for most of the range of ammonia concentrations studied, by up to 250%. This difference is too high to arise simply from enhanced electronic conductivity (due to Ni³⁺ defects⁵⁸) or better ionic conductivity and mass transport through its open structure. After all, this difference in current with ammonia is quite larger than the ∼30% difference without ammonia, where the peak current varies between 0.75 ± 0.16 mA cm⁻² for α-Ni(OH)₂, and 0.53 ± 0.15 mA cm⁻² for β-Ni(OH)₂. This reveals that the higher AOR currents on α-Ni(OH)₂ do not arise from inherent electrochemical differences between the phases, but rather due to better electrocatalytic activity of the α/γ couple. We believe that our findings can explain recent observations in the literature,⁸⁴ where a naturally oxidized nickel foam (probably consisting of amorphous α-Ni(OH)₂, as identified by earlier onsets of the NF5 and NF10 samples) exhibited higher AOR currents than the aged samples (probably consisting of β-Ni(OH)₂, as identified by belated onsets of the NF50 sample). Interestingly, at ammonia concentrations above 1 M, the AOR current density starts to decay. This will be discussed in the next section.

To investigate the effect of pH on AOR, the two catalysts were cycled with 0.5 M ammonia, and the pH was adjusted in the range 11 to 13 (Fig. 5b). This range was chosen to produce the electrochemically active NH₃ form (rather than NH₄⁺), which is expected to be present at pH > 11 in over 98% fraction (NH₄⁺ pK_a = 9.24).^70,85 Moreover, the electro-dissolution of Ni²⁺ in the presence of ammonia, to form a [Ni(NH₃)₆]²⁺ complex, is expected to decrease at alkaline pH.^86,87 In our experiments, the oxidation currents increase with pH, on both α-Ni(OH)₂ and β-Ni(OH)₂. This is so, because the hydroxide ion is a reactant in both the Ni(OH)₂ oxidation (reaction (IV)) and the AOR (reaction (I)). Nevertheless, the α-Ni(OH)₂ maintains higher current densities than β-Ni(OH)₂, in the entire pH range. This demonstrates that the advantage of α-Ni(OH)₂ as an electrocatalyst is independent of the pH. Interestingly, the voltametric waves change their shapes with ammonia concentration and pH, depending on the Ni(OH)₂ phases (Fig. S2 and S3, ESI†).

As ammonia is added, the CV shows: (i) a delay in the onset potential for oxidation, (ii) a rise in anodic current, especially in higher overpotentials, and (iii) a decline in the cathodic peak current in the back scan. We quantitatively analyzed these changes using three parameters (Fig. 6): Δi to mark the change in current at E = 1.535 V vs. RHE; ΔE_onset to mark the delay in oxidation onset potential, relative to the ammonia-free electrolyte; and i_c,NH3/i_c,none (%) to mark the relative change in the cathodic back peak, when comparing electrolytes with and without NH₃. These changes are reported in comparison to the redox peaks without ammonia, to exclude any differences in conductivity. The additional current at higher ammonia concentrations (Δi, Fig. 6b) is initially higher in α-Ni(OH)₂ than in β-Ni(OH)₂. However, when [NH₃] surpasses 0.75 M, Δi plummets in both, even becoming negative in α-Ni(OH)₂. The decline in Δi in β-Ni(OH)₂ is smaller and only starts above 1 M NH₃. These observations suggests that NH₃ saturates and blocks the surface both towards the AOR (reaction (I)), and then even towards the OH⁻ needed to oxidize Ni(OH)₂ to NiOOH (reaction (VI)). This is in contrast to the UOR, where rising urea concentrations give rise to surface saturation, but no decline in the current even at 2 M urea.⁸² Moreover, such a delay in AOR onset on Ni(OH)₂ was recently observed with rising [NH₃].⁸⁴

Fig. 6 Analysis of the changes in the voltametric waves. (a) Definitions of change parameters (see text for further explanations), comparing cycle 10 without NH₃ and cycle 1 with 1 M NH₃, on α-Ni(OH)₂, in 1 M KOH + 0.5 M K₂SO₄. (b)–(d) The effect of ammonia concentration on the (b) additional oxidation current at 1.535 V, (c) delay in onset potential, and (d) the relative change in back reduction current, for α-Ni(OH)₂ (red circles) and β-Ni(OH)₂ (blue squares). (e)–(g) The effect of pH on the same parameters.

The onset potential for oxidation (as defined by the intersection of tangent lines, Fig. 6a) is delayed on both phases, even in the lowest concentrations (ΔE_onset > 0, Fig. 6c). This is in contrast to other reaction systems, where the catalytic oxidation either coincides with the Ni^II/Ni^III transition potential (like in the UOR⁸⁸) or follows after it (like in OER⁸⁹) – but does not affect the Ni^II/Ni^III redox potential itself. Here, ΔE_onset is quite constant on both phases in low ammonia concentrations, increasing dramatically at [NH₃] > 0.75 M. The delay in onset probably arises from binding of NH_x species to Ni(OH)₂, stronger here than the binding of urea, for example (where the effect is not observed). In the literature, the adsorption energies of NH_x species on Ni(OH)₂ were calculated for the case of gaseous ammonia or urea on the (0001) facet of β-Ni(OH)₂, showing comparable values around 310–340 kJ mol⁻¹.^70,90 However, calculations on the surface of α-Ni(OH)₂ are challenging, due to the ill-defined structure of this phase. When the ammonia concentration is further increased, the OH⁻ can no longer react with the Ni(OH)₂ surface to oxidize it to NiOOH (reaction (VI)), thus delaying the Ni^II/Ni^III transition.

The suppression of the transition can be further studied by analyzing the cathodic peak variation (i_c,NH3/i_c,none, Fig. 6d). Note that i_c represents only the Ni^III → Ni^II reduction, since the ammonia oxidation is irreversible. The value of i_c,NH3/i_c,none initially increases on both phases, then decreasing when [NH₃] > 0.5 M. Notably, the decline in cathodic currents at higher ammonia concentrations is more steep for α-Ni(OH)₂ than β-Ni(OH)₂, with almost an entirely irreversible process at 2 M. The decline in i_c,NH3/i_c,none after [NH₃] > 0.5 M indicates that there is less and less NiOOH available to be reduced, as the NH₃ concentration rises. This could happen if NH₃ molecules block OH⁻ from accessing surface Ni^II sites, thus generating less NiOOH to begin with. Alternatively, the NiOOH could be blocked by NH₃ molecules, if the AOR is slow enough, decreasing the effective surface available for the back reduction. Overall, the surface of Ni(OH)₂ and/or NiOOH could be blocked by NH_x species (x = 0–3), effectively poisoned similarly to poisoning of Pt by N_ads during AOR electrocatalysis.⁶² In fact, ammonia adsorption by breaking Ni–O bonds, followed by dehydrogenation, were proposed as the first AOR steps.⁷⁰

Interestingly, a decline in the cathodic current following urea oxidation on Ni(OH)₂/NiOOH was previously proposed as evidence for an indirect AOR mechanism (reaction (IV)). The reason is that the chemical reduction of Ni^III back to Ni^II by ammonia would leave less NiOOH to undergo electrochemical reduction in the back scan, resulting in a lower cathodic current.⁸² However, this is hardly the case here: the decline in AOR at higher [NH₃] – which should revive the cathodic currents, as less NiOOH is chemically back-reduced – actually leads to lower, not higher, cathodic currents. Thus, a mechanism switch is hard to confirm here, and surface poisoning remains the most viable hypothesis. Furthermore, one must remember that the Ni(OH)₂ electrodes can contain a mosaic of phases that are ever changing at the timescale of the CV experiment, complicating a possible discussion of a mechanism shift.

Since a competition between ammonia and hydroxide ions seems to play a key role in the observed trends, Δi, ΔE_onset and then changes to the cathodic back-reduction peak were analyzed for α-Ni(OH)₂ and β-Ni(OH)₂ at different electrolyte pH (Fig. 6e, f and g). The ammonia concentration was set to 0.5 M, to maximize the difference between the phases. Interestingly, the effect of pH was found to be much weaker than that of the ammonia concentration. There is a small increase in Δi with pH; since OH⁻ is a reactant for both the nickel hydroxide oxidation and the AOR (reactions (VI) and (I)), the current is indeed expected to rise with increasing concentration of hydroxide ions. The near-constant ΔE_onset at pH > 11, suggests that the oxidation is no longer limited by [OH⁻], indicating that surface blocking by NH₃ becomes key. Note that by defining the onset delay as the reviewed parameter, rather than the onset potential, we can focus solely on the differences originating from the NH₃/OH⁻ ratios on the pH-insensitive RHE.

The value of i_c,NH3/i_c,none is somewhat higher at lower pH, reaching higher than 100% values. This could be due to some electrodissolution of Ni in ammonia to form nickel ammonium complexes, which intensifies in lower pH;^86,87 the reverse redeposition of Ni could be the side reduction process. However, dissolution (at open circuit) is minimal (see below). Nickel hydroxide can extract Fe traces from the electrolyte, leading to iron doping which affects its electrochemistry,^74,91 including for electrocatalysis of the UOR⁹² and the OER.⁹³ To rule out the effect of Fe impurities on the observed trends in AOR, we repeated the CV cycling in Fe-free solutions at several ammonia concentrations (Fig. S6 and S7, ESI†). The resulting voltammograms present the same trends as observed in Fig. 6 for Δi and ΔE_onset, differing somewhat in the current densities, and showing a 30 mV earlier onset for α-Ni(OH)₂. Hence, the trends in onset delays and current variations vs. [NH₃] are unrelated to any iron incorporated in the Ni(OH)₂.

Surface poisoning by ammonia

To investigate the poisoning of Ni(OH)₂/NiOOH surfaces in high ammonia concentrations, the electrodes were either soaked or voltammetrically cycled in 2 M ammonia, and then transferred to an ammonia-free electrolyte. If the ammonia is weakly adsorbed to the surface, it should desorb (at least partially) in the NH₃-free electrolyte. Then, the hydroxide ions would again be able to oxidize the surface and the nickel peak current density would recover. However, the nickel redox peaks are never fully restored after switching to the ammonia-free electrolyte, for neither α-Ni(OH)₂ or β-Ni(OH)₂ (Fig. 7 and Fig. S8, ESI†). Notably, the surface adsorption is stronger on α-Ni(OH)₂: the currents do not recover at all, and the Ni peaks are almost eliminated. On β-Ni(OH)₂, most of the current is restored after soaking in NH₃ (Fig. 7b), but less is restored after CV cycling (Fig. S8, ESI†). Overall, these observations suggest that α-Ni(OH)₂ binds NH₃ (or a NH_x intermediate) more strongly than β-Ni(OH)₂, and the surface remains poisoned even in the ammonia-free electrolyte. To gauge the possibility of Ni electro-dissolution, which could be enhanced by the presence of NH₃,^86,87 we measured the Ni content in the electrolyte by ICP-MS following soaking in 1 mL of 2 M ammonia solution for 10 minutes, as the duration of an electrochemical measurement in ammonia. The results show that only about 1% of the α-Ni(OH)₂ electrode mass has been dissolved. Therefore, the dissolution of nickel does not explain the huge decay in activity after soaking in ammonia solution, appearing at Fig. 7a.

Fig. 7 CV of nickel hydroxide in an ammonia free solution before (10th cycle, in black) and after (2nd cycle, blue) soaking the electrode in 2 M NH₃ solution for 10 minutes (same as the duration of an AOR CV experiment) for (a) α-Ni(OH)₂ and (b) β-Ni(OH)₂.

The difference in NH₃ adsorption energy between α-Ni(OH)₂ and β-Ni(OH)₂ corroborates the trends in AOR electrocatalysis (Fig. 8). At low NH₃ concentrations, the AOR currents are higher on α-Ni(OH)₂ since it adsorbs NH₃ more strongly, promoting dehydrogenation steps. Meanwhile, at high NH₃ concentrations, strongly adsorbed ammonia poisons the sites and decreases the local pH, thus causing large onset delays and the drop in peak current. Furthermore, the fact that NH₃ adsorbs strongly on Ni(OH)₂ explains the delay in the onset potential for oxidation to NiOOH, which was first observed here. To the best of our knowledge, this is the first time that irreversible poisoning of α-Ni(OH)₂ by ammonia is reported. The weaker adsorption to β-Ni(OH)₂ explains the lower currents and the low dependance on ammonia concentrations for the trends of Δi and ΔE in Fig. 6. Moreover, reexamining the RRDE experiment in Fig. 4, we notice that the AOR currents almost disappear at high ammonia concentrations, despite the rotation eliminating mass transfer limitations. Then, at very high potentials (E > 1.75 V vs. RHE), the current increases, while O₂ is not produced, according to the ring currents. Possibly, the strongly adsorbed NH_x species are getting released by forming NO_x species, similarly to N_ads on Pt.^62,94

Fig. 8 Schematic of the AOR on Ni(OH)₂/NiOOH, comparing the α/γ vs. the β/β couple as a function of ammonia concentration. Ammonia binds more strongly to α-Ni(OH)₂, leading to better electrocatalysis at intermediate [NH₃], and to irreversible poisoning at high [NH₃].

Interestingly, the poisoning related phenomena are more pronounced for α-Ni(OH)₂ than for β-Ni(OH)₂, i.e. the current decline and the onset delay at ammonia saturation. This further suggests a stronger adsorption energy of ammonia for the surface of α-Ni(OH)₂ than β-Ni(OH)₂. This is supported by how the currents can be recovered on β-Ni(OH)₂ in the desorption experiments (Fig. 7), while α-Ni(OH)₂ remains blocked.

Could the different responses of α-Ni(OH)₂ and β-Ni(OH)₂ to surface poisoning arise from different electrochemical surface areas (ECSAs)? Unfortunately, the ECSA of nickel hydroxide phases is hard to determine, due to interfering contributions from intercalation, overcharging, and pseudocapacitance.⁷⁵ Herein, we measured the ECSA by two methods: (1) double layer capacitance measured by CV in the non-faradaic region, and (2) electrochemical impedance spectroscopy (EIS), as proposed by Watzele et al.⁷⁶ (Fig. S9, ESI†). In the EIS method, the adsorption capacitance (C_a) of oxygenated intermediates, formed at the onset of the OER, is extracted from EIS spectra by fitting them to an equivalent circuit (Fig. S9, ESI†). At potentials where these species reach full surface coverage, this value reaches a plateau, and can be correlated to the electrochemically active surface area. The double-layer capacitance values measured by the first method were similar for the two phases: 6 μF for α-Ni(OH)₂ and 7 μF for β-Ni(OH)₂. Typically, ECSA values are calculated from the double-layer capacitance by dividing it by the specific capacitance of the material. However, this value is also debated in the literature, with reported values ranging from 40 to 300 μF cm⁻².^95,96 Therefore, we can examine the capacitance values comparatively. The EIS method yielded capacitance values of 80 μF for α-Ni(OH)₂ and 9 μF for β-Ni(OH)₂ (Fig. S9, ESI†). Each method has a priori limitations: double-layer capacitance by CV may include processes such as intercalation, adding non-faradaic currents, while the new EIS approach shows inconsistent dependency of C_a on potential, similarly to what we observe (Fig. S9, ESI†). Regardless if the ECSA of α-Ni(OH)₂ is equal or even 10 times higher than that of β-Ni(OH)₂, it is clear that it is not lower, and thus the α-Ni(OH)₂ poisoning effect is not related to surface area limitations, but rather due to stronger adsorption.

Conclusions

The two phases of nickel hydroxide, disordered α-Ni(OH)₂ and crystalline β-Ni(OH)₂, were synthesized, characterized and comparatively evaluated for their electrocatalytic performance towards the ammonia oxidation reaction. Both γ-NiOOH and β-NiOOH (resulting from oxidation of α-Ni(OH)₂ and β-Ni(OH)₂, respectively) can catalyze the AOR. Both yield N₂, as determined by differential electrochemical mass spectrometry. The α-Ni(OH)₂ electrocatalyst is more reactive towards ammonia oxidation, throughout most of the concentration range between 0.01 M and 2 M. Its catalytic activity drops at high ammonia concentrations, with lower currents and belated onset potentials when [NH₃] > 1 M. Soaking experiments confirm a surface poisoning effect, where NH₃ and/or other NH_x intermediates adsorb to the Ni(OH)₂ surface, blocking access for the OH⁻ ions required for the Ni(OH)₂ → NiOOH oxidation and for the AOR. The surface poisoning is completely irreversible for α-Ni(OH)₂, and partially reversible on β-Ni(OH)₂, suggesting stronger adsorption energies on the disordered, intercalating α-Ni(OH)₂ phase. Overall, our findings demonstrate the key role of the phases of the Ni(OH)₂/NiOOH electrocatalysts on ammonia oxidation, as well as the strong effects of poisoning and competition with hydroxide-consuming reactions. We hope that these findings will broaden our understanding of electrocatalysis of nitrogen-based fuels, and of metal oxide electrocatalysts in general.

Author contributions

I. O. P. performed the electrochemical experiments (with H. A. A. and A. H.), the DEMS experiments (with T. K. S.) and all characterizations. I. O. P. wrote the original draft. D. E. coordinated the project and edited the manuscript.

Data availability

Additional electrochemical measurements and analysis have been included in the ESI.†

Conflicts of interest

There are no conflicts to declare.

Acknowledgements

We thank the Israel Ministry of Science and Technology (grant 1001556424) and the Israel Ministry of Energy (grant 222-11-059) for partial funding, and Dr Thomas Hartman for graphic design. I. O. P. thanks the Israeli Smart Transportation Research Center (ISTRC) for a graduate fellowship. T. K. S. thanks the Azrieli Foundation and the Israel Academy of Science and Humanities for post-doctoral fellowships.

Notes and references

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Footnote

† Electronic supplementary information (ESI) available: Additional electrochemical measurements and analysis. See DOI: https://doi.org/10.1039/d4cp02950j

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