Insight into an electrodeposition process of bimetal co-substituted α-Ni(OH)₂ preparation

2.1. Preparation of the electrodeposition system

Analytical−grade NiCl₂·6H₂O, AlCl₃·6H₂O, MnCl₂, NaCl, and anhydrous ethanol were purchased from Sinopharm Group Chemical Reagent Co., LTD, and used without further purification. All electrodepositions were carried out in a three−chamber electrochemical cell, featuring two symmetrically arranged anode chambers flanking the cathode chamber. The compartments were separated by a cation−exchange membrane (CEM, HTECH-18231123, HUAMO TECH, China) with selective cation permeability. The cathode was a 304 stainless-steel sheet (35 mm × 90 mm × 1 mm, LHW), while the anode was an iridium-ruthenium-coated titanium sheet (17.5 mm × 90 mm × 1 mm, LHW).

Previous studies [35] have shown that adding ethanol to the catholyte reduces pH fluctuations during electrolysis. Moreover, ethanol can form coordination complexes with Ni²⁺, increasing the overpotential for Ni²⁺ reduction to metallic nickel and thereby suppressing undesired side reactions. However, a higher ethanol content increases cell voltage; therefore, a 50% ethanol−water catholyte and deionized water as the anolyte were used.

In reported α-Ni(OH)₂ containing substituted metals, the typical molar ratio of Al:Ni is ∼0.1:1, whereas Mn:Ni is generally <0.1:1 [22,25]. To obtain samples with Mn:Al:Ni molar ratios comparable to literature, electrolytes (x₁ M MnCl₂ + 0.02 M AlCl₃ + 0.18 M NiCl₂, x₁ = 0.0036–0.018) were prepared to yield Mn:Al:Ni ratios of 0.02:0.11:1, 0.04:0.11:1, 0.06:0.11:1, 0.08:0.11:1, and 0.1:0.11:1 for both catholyte and anolyte. The corresponding deposition system is referred to as Experiment I. Samples from Experiment I were designated Ni−Al−Mn_+-−y₁ (_+- indicates the presence of MnCl₂ in both catholyte and anolyte; y₁ is the initial MnCl₂ molar ratio, y₁ = 0.02–0.1). For Experiment II, the anolyte composition remained the same as in Experiment I, whereas MnCl₂ was removed from the catholyte. Samples were labeled Ni−Al−Mn₊−y₂ (₊ indicates MnCl₂ present only in the anolyte; y₂ = y₁ = 0.02 – 0.1). In Experiment III, the catholyte of Experiment II was retained, and the MnCl₂ concentration in the anolyte was increased tenfold. These samples were labeled Ni−Al−Mn₊−y₃ (₊ denotes MnCl₂ only in the anolyte; y3 = 10 y2 = 0.2 – 1.0). Data from an electrodeposition system containing 0.02 M AlCl₃ + 0.18 M NiCl₂ as both anolyte and catholyte were also considered, designated as Experiment 0, with the resulting sample labeled Ni-Al.

The corresponding experimental groups, electrolyte compositions, and sample labels are summarized in Table 1.

2.2. Measurements of the electrodeposition systems

The pH of the catholyte was measured using a pH meter (LEICI, PHS−3E, China). The electrode was kept immersed in the catholyte throughout each experiment until power termination. As shown in Figure 1, the pH probe was fixed 25 mm away from the cathode sheet.

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Figure 1.

Diagram of the relative position between the pH electrode and the cathode sheet.

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To clarify the main electrochemical reactions involved in the electrodeposition process, cyclic voltammetry (CV) was carried out using a self-assembled three-electrode configuration on an electrochemical workstation (Zennium−pro, Zahner PP211, Germany). The working (D = 1 mm) and counter (D = 1 mm) electrodes were fabricated from the same materials as the cathode and anode used in electrodeposition. All CV measurements were referenced to an Ag/AgCl electrode and recorded in the cathodic direction, followed by the anodic sweep at a scan rate of 30 mV/s. NiCl₂, AlCl₃–NiCl₂, and MnCl₂–AlCl₃–NiCl₂ solutions prepared with a water-ethanol solvent (1:1, v/v) were used as electrolytes for CV, and NaCl was added to each system to increase conductivity. The concentration of all solutes, including NaCl, was maintained at 1 g/L.

2.3. Measurements of the samples

An inductively coupled plasma atomic emission spectrometer (ICP, Leeman Prodigy XP, USA) was used to determine the metal ion content of the samples. X-ray diffraction (XRD) was performed on a diffractometer (Bruker D8, USA) operated at 40 kV and 40 mA using CuKα radiation (λ = 1.54 Å) to identify the crystalline phase. Fourier-transform infrared spectroscopy (FT-IR) was carried out on an infrared spectrometer (Nicolet iS50, USA) to further assist in phase analysis.

To assess alkaline stability, 1 g of each sample was immersed in 100 mL of 6 M KOH solution. At scheduled intervals, samples were withdrawn and subjected to XRD to monitor phase evolution over time.

3.1. Cyclic voltammetry studies of the electrodeposition systems

CV curves related to the electrodepositions mentioned are presented in Figure 2. Since the cathodic reaction is effective for electrodeposition, scans in the negative potential direction are selected for analysis.

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Figure 2.

CV curves for various electrodeposition systems.

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According to the literature [33], the electrochemical reactions potentially occurring at the cathode during electrodeposition are presented in Eqs. (1–3). Among these, Eq. (1) is primarily responsible for α-Ni(OH)₂ synthesis, whereas Eqs. (2, 3) are considered side reactions. On the basis of the standard electrode potentials (E^Θ) of these reactions, Eq. (3) is the least likely to occur owing to its unfavorable thermodynamics. The similar E^Θ values of Eqs. (1, 2) suggest their possible simultaneous occurrence during electrodeposition.

However, previous studies have shown that ethanol in a catholyte can effectively suppress the occurrence of Eq. (3) [35]. Therefore, irrespective of the electrolyte component, all the CV curves have the same shape, and only one reaction can be captured, which has been proven to be Eq. (1) [36]. Moreover, no reaction with respect to metal deposition was observed, indicating that no metal byproduct was produced in the electrodeposition systems involved in the present work.

3.2. Analysis of the electrodeposition process

The molar ratios of Mn:Al:Ni for the samples from Experiment І (Ni−Al−Mn_+-−y₁) are tabulated in Table 2. For comparison, data corresponding to electrodeposition in an electrolyte with an Al:Ni molar ratio of 0.1:1 (Experiment 0) are also included in Table 1. R_Mn and R_Al were separately used to intuitively represent the ratio of the proportion of Mn and Al in the sample to that in the initial electrolyte. The R_Mn and R_Al for Experiment І are denoted as R_Mn^І and R_Al^І, respectively.

As shown in Table 2, the R_Mn^І values are consistently close to 1, indicating that the Mn molar ratio in the samples from Experiment І is close to that in the initial electrolyte.

With respect to the Al ratio, our previous findings demonstrated a slight enrichment of Al in samples compared with the initial electrolyte during electrodeposition from an AlCl₃−NiCl₂ solution [34]. As shown in Table 2, a R_Al value of 1.09 was obtained for the electrodeposition conducted in 0.02 M AlCl₃ + 0.18 M NiCl₂ solution (Experiment 0).

The R_Al values for Experiment I are consistently near 2.5, demonstrating that the Al fraction in the samples is about 2.5−fold higher than that in the initial electrolyte. This reflects a pronounced deviation in Al incorporation relative to its initial concentration.

Since the concentrations of NiCl₂ and AlCl₃ were identical in Experiment 0 and Experiment I, the presence of MnCl₂ is identified as the cause of this divergence. Specifically, Mn²⁺ in the electrolyte appears to modulate the deposition behavior of Al³⁺.

Because deposition occurs exclusively in the cathodic chamber, a direct method to reduce Al enrichment in the sample would be to lower MnCl₂ content in the catholyte. However, as indicated in Table 2, even at a Mn:Ni ratio as low as 0.02:1, a significant deviation persists, demonstrating that reducing MnCl₂ concentration in the catholyte does not effectively mitigate the observed excess Al incorporation.

Therefore, an extreme approach was adopted by directly removing MnCl₂ from the catholyte. In this configuration, x₁ M MnCl₂ + 0.02 M AlCl₃ + 0.18 M NiCl₂ solution served as the anolyte, while 0.02 M AlCl₃ + 0.18 M NiCl₂ acted as the catholyte (Experiment II). The Mn:Al:Ni molar ratios of the resulting samples (Ni−Al−Mn₊−y₂) are listed in Table 3.

As shown in Table 3, the R_Al^Ⅱ values are about 1.15, which is very close to 1.09 observed in Experiment 0. This indicates that introducing MnCl₂ exclusively into the anolyte effectively suppresses the influence of Mn²⁺ on the incorporation of Al³⁺.

In contrast, the R_Mn^II value is about 0.1, meaning that the Mn fraction in the resultant samples is only one-tenth of that in the initial anolyte. Since the catholyte in Experiment II is MnCl₂-free, the Mn present in the samples originates solely from diffusion and electromigration from the anolyte. The amount of Mn²⁺ transported into the catholyte is therefore much lower than its original concentration, resulting in a significantly reduced Mn ratio in the product.

To achieve Mn:Al:Ni molar ratios comparable to those reported in the literature [24,25], Experiment III was carried out with the MnCl₂ concentration in the anolyte increased tenfold relative to Experiment II. Table 4 summarizes the Mn:Al:Ni ratios in both the anolyte and the corresponding samples (Ni−Al−Mn₊−y₃, y₃ = 0.2 − 1.0).

As seen from Table 4, the Mn content in the samples of Experiment III remains about one-tenth of that in the initial anolyte. However, because the anolyte MnCl₂ concentration was increased tenfold relative to Experiment II, the Mn fraction in the deposited samples increased by a similar magnitude. This confirms that elevating Mn content in the anolyte results in a proportional rise in Mn content in the sample.

With regard to the Al ratio, R_Al increased slightly from ∼1.15 in Experiment II to ∼1.22 in Experiment III, mainly due to the greater influence of Mn²⁺ on the deposition behavior of Al³⁺ at higher Mn²⁺ concentration.

A clear trend can therefore be drawn from Tables (2-4), which is also reflected in the data presented in Table 5.

Table 5 shows that samples with comparable Mn:Ni ratios exhibit substantial differences in their Al:Ni ratios between the two experiments.Throughout all experiments, both the applied current and the electrode surface area in contact with the electrolyte remained constant. Meanwhile, the natural pH of the electrolytes (ranging from 3.3–3.5) displayed only minor fluctuations. Therefore, the rate of OH^– generation at the cathode surface is reasonably assumed to be nearly identical among all conditions.

The Ksp values of Ni(OH)₂, Al(OH)₃, and Mn(OH)₂ are 5.5 × 10^–16, 1.3 × 10^–33, and 2.1 × 10^–13, respectively. Accordingly, the pH values required for the precipitation of Ni²⁺, Al³⁺, and Mn²⁺ can be calculated from Eqs. (4–6).

In the chemical precipitation of α-Ni(OH)₂ containing substituent ions, the overall precipitation process is generally considered to proceed via co-precipitation [14,23], as represented in reaction (7).

Regarding the electrodeposition route, two mechanistic pathways have been described in the literature for the formation of α-Ni(OH)₂ containing Al substitution [32,33,37,38].

The first pathway, described by reactions (8–9), involves the initial precipitation of Al(OH)₃ as the OH^– concentration increases near the cathode, followed by its dissolution and participation in the formation of a hydrotalcite-like phase.

The second pathway, shown in reaction (10), involves the direct co-precipitation of Ni²⁺, Al³⁺, and other substituent ions with OH^– to form α-Ni(OH)₂−type phases.

Furthermore, the hypotheses and inferences in the literature suggest that a lower OH^– generation rate favors the first pathway, while a higher rate promotes the second pathway [28,32].

Based on the observations of the present work, the reactions most likely follow the second (co-precipitation) pathway. Supporting evidence is provided below.

Figure 3 shows photographs of the cathode chamber and cathode plate after electrodeposition in Experiment III. As seen, the product adhered firmly to the cathode sheet with no observable sediment accumulation at the bottom of the chamber.

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Figure 3.

Cathode chamber and cathode plate at the end of electrodeposition for various anolyte compositions (Experiment III): (a) Mn:Al:Ni = 0.2:0.11:1, (b) Mn:Al:Ni = 0.4:0.11:1, (c) Mn:Al:Ni = 0.6:0.11:1, (d) Mn:Al:Ni = 0.8:0.11:1, (e) Mn:Al:Ni = 1.0:0.11:1.

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Meanwhile, when the maximum concentration of Al³⁺ (0.02 M) is substituted into Eq. (5), the minimum pH at which Al³⁺ precipitation initiates in the catholyte is calculated to be 3.6. However, as shown in Figure 4, the pH of the catholyte fluctuated between 3.25 and 3.7 during electrodeposition. Therefore, the bulk solution does not satisfy the thermodynamic conditions for precipitation, confirming that precipitation does not occur in the bulk phase. Electrodeposition in which the product forms directly on the electrode surface is referred to as in-situ electrodeposition [39,40], and such in-situ behavior was observed in all experiments of this study. According to previous reports, this characteristic is associated with Al substitution [28].

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Figure 4.

The catholyte pH during electrodeposition for Experiment Ⅲ.

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Once electrodeposition begins, OH^– is continuously generated at the cathode−catholyte interface, resulting in persistently high local OH^– concentration near the electrode. Thus, throughout the in-situ deposition process, product formation occurs in a locally high−pH microenvironment. The precipitation behavior of the metal ions involved can therefore be regarded as consistent with co-precipitation following reaction (10). The co-precipitation pathway applicable to this work is summarized below.

(І) Firstly, with power applied, the cathode−catholyte interface is rapidly enriched with OH^–. The nearby cations coordinate with OH^–, leading to a decrease in both interfacial metal-ion concentration and local pH. In multi−component electrolytes, different metal ions first undergo hydrolytic speciation before precipitating. For instance, Ni²⁺ can form intermediates such as NiOH⁺ and Ni₄(OH)₄⁴⁺ via reaction (11–12) [37,41], while Mn²⁺ may form MnOH⁺ through reaction (13) [42]. Due to the strongly alkaline interfacial environment, Al³⁺ is converted directly to AlO₂^– as depicted in reaction (14).

(Ⅱ) Secondly, driven by ‌concentration gradients‌ and/or ‌electrostatic forces‌, the intermediates and AlO₂^– then diffuse or migrate toward the bulk solution, while metal cations from the bulk move toward the interface to replenish those consumed. Meanwhile, continued electrolysis maintains the interfacial OH^– supply, raising the local pH. As a result, metal cations, intermediates, AlO₂^–, and OH^– coexist in the vicinity of the cathode, where interfacial precipitation gradually proceeds.

(Ⅲ) Thirdly, as long as metal ions and OH^– are continuously replenished at the interface, their concentrations eventually reach the threshold for co-precipitation, yielding α-Ni(OH)₂ containing substituted metal atoms. Owing to the ‌dynamic balance‌ between the consumption and replenishment of these species at the interface, the concentrations of metal ions, intermediates, and OH⁻ ions remain within the ‌precipitation-suitable range‌, enabling the continuous synthesis of α-Ni(OH)₂ with substituted metal atoms.

For reaction (10), the Ksp of the product should be between 1.3×10^–33 (Al(OH)₃) and 2.1×10 ^–13 (Mn(OH)₂). According to Eqs. (4–6), the minimum pH at which co-precipitation begins is between that for Al(OH)₃ (3.6) and that for Mn(OH)₂ (12.77). Consequently, a relatively lower pH‌ in the catholyte near the interface during co-precipitation promotes a ‌higher Al content‌ in the product, whereas a ‌higher pH‌ results in a ‌lower Al content‌.

In the electrolyte system where both anolyte and catholyte consist of AlCl₃–NiCl₂ (Experiment 0, Figure 5a), Ni²⁺ and Al³⁺ primarily regulate the interfacial pH by consuming OH^– during electrolysis, leading to a slight enrichment of Al in the final product (R_Al = 1.09).

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Figure 5.

The distribution of metal-ion species in the electrodeposition systems: (a) Experiment 0, (b) Experiment І, (c) Experiment II, and (d) Experiments Ⅲ.

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In MnCl₂-containing systems (Experiments І–III, Figures 5b-d), Mn²⁺ and its intermediates also participate in OH^– consumption. Since the OH^– generation rate is constant for all systems, the presence of Mn²⁺ lowers the interfacial pH relative to the Mn-free system (Figure 5a). A lower local pH is more favorable for Al³⁺ precipitation, explaining the consistently higher Al:Ni ratios in MnCl₂-containing samples compared to Mn-free systems, as shown in Tables (2-4).

In Experiment I, MnCl₂–AlCl₃–NiCl₂ served as both anolyte and catholyte (Figure 5b). The Al:Ni ratio in the samples was markedly higher than that obtained from the Mn-free system; however, no clear upward trend in Al content was observed with increasing MnCl₂ concentration, likely because the interfacial availability of Al³⁺ and its hydrolyzed species becomes the limiting factor once co-precipitation is triggered.

Compared with Experiment I, MnCl₂–AlCl₃–NiCl₂ solution served only as the anolyte in Experiment II. All Mn²⁺ ions entering the cathode chamber were therefore introduced exclusively via diffusion and electromigration from the anolyte (Figure 5c). During co-precipitation, the concentration of Mn²⁺ and its hydrolytic intermediates at the interface was lower than in Experiment I, resulting in a weaker pH−buffering effect. Consequently, the Al:Ni ratio in samples from Experiment II was noticeably lower than that from Experiment I.

Experiment III comprised a series of systems in which the MnCl₂ concentration in the anolyte was increased tenfold relative to Experiment II (Figure 5d). This enhanced binding capacity for OH⁻ ions led to a lower interface pH. Consequently, the Al ratio in the samples obtained from Experiment III was higher than that from Experiment II.

Notably, as shown in Table 5, the samples with nearly identical Mn:Ni ratios obtained from Experiments І and Ⅲ clearly deviate in the Al:Ni ratio. In the experiments described in this paper, the generation rate of OH^– was considered to be constant, and the initial concentrations of NiCl₂ and AlCl₃ in the catholyte and anolyte were identical. Therefore, the significant difference in the Al:Ni ratios of the samples indicates that there is likely to be an obvious difference in the pH at the interface. The higher Al ratios in the samples from Experiment I indicate a lower pH at the corresponding interface. Furthermore, more OH^– might be consumed by Mn²⁺ and its intermediates in Experiment І with a MnCl₂−AlCl₃−NiCl₂ solution as the catholyte and anolyte than in Experiment Ⅲ with a MnCl₂−AlCl₃−NiCl₂ solution as the anolyte and an AlCl₃−NiCl₂ solution as the catholyte. The almost identical Mn:Ni ratios for the samples from Experiments І and Ⅲ might be because more Mn intermediates migrated to the bulk solution, driven by concentration gradients‌ and/or electrostatic forces before co-precipitation in Experiment І. The two reasons above result in the concentration of Mn²⁺ and its intermediates that can take part in the co-precipitation process for Experiment І being comparable to that for Experiment Ⅲ, leading to a comparable Mn:Ni ratio in the final samples.

3.3 Characterization of the samples

In metal-substituted α-Ni(OH)₂, substitution involves replacing a portion of Ni atoms in the α-Ni(OH)₂ lattice with other metal atoms [15], resulting in structural characteristics consistent with those of pure α-Ni(OH)₂ [43]. α-Ni(OH)₂ consists of layered Ni(OH)₂ sheets. Because there are insufficient OH^– ions within the lattice, to neutralize the positive charge of these layers [4], interlayer water molecules and anions are inserted to balance charge [16]. The irregular size and random distribution of these intercalated groups disrupt periodic stacking, producing a disordered interlayer structure [44]. Consequently, XRD patterns typically display broadened and asymmetric diffraction peaks indicative of structural disorder [8,12].

As shown in Figure 6(a), the asymmetric diffraction peaks in all the XRD patterns indicate the typical chaotic internal structure of α-Ni(OH)₂. The diffraction peaks located at approximately 12°, 23°, 33°, and 60° can be indexed to the (003), (006), (101), and (110) planes for α-Ni(OH)₂. Compared with the sample (Ni-Al) from Experiment 0, the (015) plane for α-Ni(OH)₂ disappears from the XRD patterns for the samples obtained from Experiment III. In addition to the abovementioned diffraction peaks, no diffraction peaks belonging to other precipitates are captured in the XRD patterns in Figure 6(a), indicating that no other precipitates were formed except for α-Ni(OH)₂ and that Al or Mn atoms were successfully embedded into the lattice of α-Ni(OH)₂.

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Figure 6.

(a) The XRD patterns and (b) FT-IR spectra for the samples obtained from Experiment Ⅲ.

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To complement the phase identification via XRD, FT-IR analysis was performed (Figure 6b). All the spectra exhibit identical profiles. The wide band at approximately 3460 cm^–1 and the weak band at about 1640 cm^–1 correspond to the vibration of the H-bonded OH^– groups and the bending vibration of the H₂O molecules, respectively [9,45], which provides evidence of the existence of intercalated H₂O molecules [7]. The tiny peak at about 1380 cm^–1 is related to the asymmetric vibration of the interlayer CO₃^2– [18,20]. The symmetric vibration of CO₃^2– at about 1480 cm^–1 is not observed in the spectra of the samples from Experiment III, which may be combined with the nearby asymmetric vibration of CO₃^2–. The vibration of the in-plane Ni-O-H peak at approximately 640 cm^–1 moved in the low wavenumber direction for the samples from Experiment III, which may be the result of the combined action of multiple metal elements of Ni/Al/Mn in the lattice. From the analysis above, all the samples mentioned above have a layered structure with a Ni-O-H layer, and there are water molecules as well as other anions in the interspaces, which are structural characteristics of α-Ni(OH)₂ [45,46]. Furthermore, the analysis of the samples from Experiment III demonstrated that the Al and Mn atoms were successfully embedded into the α-Ni(OH)₂ lattice.

To evaluate alkaline stability, α-Ni(OH)₂ samples are commonly immersed in 6 M KOH and periodically analyzed to monitor phase evolution before and after prolonged soaking [26,44]. This method was likewise employed in the present study to assess the alkaline stability of the synthesized materials.

Previous studies have shown that α-Ni(OH)₂ synthesized from a pure NiCl₂ electrolyte can be completely transformed into β-Ni(OH)₂ after only one day of immersion in strong alkali [34]. To assess the stability of the samples from Experiment III, they were immersed in 6 M KOH and periodically collected for phase identification. The results are presented in Figure 7.

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Figure 7.

XRD patterns of different samples soaked in 6 M KOH solution for various durations: (a) Ni−Al, (b) Ni−Al−Mn₊–0.2 to Ni−Al−Mn₊–1.0, and (c) Ni−Al−Mn₊–1.0 compared with the sample from the literature.

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As illustrated in Figure 7(a), the XRD patterns of the Al-substituted α-Ni(OH)₂ sample still display most α-Ni(OH)₂ diffraction peaks after 5 days of immersion; however, peaks characteristic of β-Ni(OH)₂ begin to emerge. In contrast, for the Al–Mn co-substituted sample from Experiment III (Figure 7b), all diffraction peaks associated with α-Ni(OH)₂ persist even after 30 days of immersion. Except for a peak between 61° and 63°, attributed to the overlap of α-Ni(OH)₂ and β-Ni(OH)₂, no β-Ni(OH)₂ reflections are observed. This confirms that the stability of α-Ni(OH)₂ is strengthened by Al substitution and further enhanced by Al–Mn co-substitution.

To further compare stability, two samples with similar Mn:Al:Ni ratios—one prepared by electrodeposition (0.0987:0.133:1) and another by chemical precipitation (0.1:0.144:1)—were immersed in 6 M KOH for 30 days. The corresponding XRD patterns are shown in Figure 7(c).

After immersion, the diffraction peak at ∼12° (assigned to α-Ni(OH)₂) is still retained in the electrodeposited sample but splits into two distinct peaks in the chemically precipitated sample. The α-Ni(OH)₂ reflections near 23°, 33°, and 39° in the precipitated sample shift toward lower diffraction angles, positioning the peaks at 33° and 39° to the left of the β-Ni(OH)₂ reflections at similar positions. Furthermore, the α-Ni(OH)₂ peak at ∼23° in the chemically precipitated sample lies closer to the β-Ni(OH)₂ peak at ∼18°.

Both samples retain the α-Ni(OH)₂ diffraction band near 60°, and a mixed α-/β-Ni(OH)₂ signal appears near 51°. A new reflection emerges near 62° in both spectra, but in the chemically precipitated sample this peak is located closer to the β-Ni(OH)₂ position.

These observations collectively indicate that the Al–Mn co-substituted α-Ni(OH)₂ synthesized via electrodeposition exhibits superior alkaline stability compared to its chemically precipitated counterpart.