Exploring the Voltage Stability of Birnessitic Bi-MnO₂ Cathodes Formed In Situ in NaOH and KOH-Based Electrolytes

All electrochemical tests were done using Hg/HgO reference electrodes filled with 20 wt% KOH. Two test fixtures were used in this study. The first was a flooded beaker cell, illustrated in SI Fig. 1. The beaker cell was comprised of a polypropylene bottle, a 6.6 M or 10.5 M electrolyte (NaOH or KOH) and a platinum or nickel foam counter electrode. 6.6 M corresponds to 21 and 37 wt% NaOH and KOH, respectively.

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The second fixture was an electrolyte-lean T-cell, illustrated in SI Fig. 2, composed of a polypropylene tee-junction with a 16 mm diameter. In the center of the tee, the working electrode, two synthetic cotton separators, and the counter electrode were sandwiched between two graphite current collectors and two 304 stainless steel plungers. Approximately 1–2 ml of an aqueous electrolyte was added into the system. NiOOH counter electrodes were prepared by mixing NiOOH and PTFE with mortar and pestle in a 9:1 mass ratio and pressing the mixture into nickel mesh. NiOOH was prepared by dissolving 1.43 g Na₂S₂O₈ in 100 ml DI water, then adding 1.41 g Ni(NO₃)₂·6H₂O. Upon dissolution, enough KOH was added to make the solution a strong base (pH > 13), resulting in the formation of black precipitate. The mixture was stirred at 45 C for 2 h, then filtered to extract the NiOOH powder.

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Following the electrochemical synthesis method reported by Yadav et al., ² each cathode was discharged in 6.6 M KOH or NaOH at a C/3 rate to −1.0 V, allowed to rest for 5 min, and then charged at a C/3 rate to 0.3 V. The resulting voltage-capacity curve in KOH was as expected and is shown in SI Fig. 3, annotated with the expected reaction pathway (as reported by Yadav et al. for KOH). ²

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Bi-birnessite cathodes were also cycled in 6.6 M KOH or NaOH using the cyclic voltammetry (CV) technique on the Biologic VMP3 after undergoing the birnessite formation process. The CV method used applied a voltage sweep at a rate of 0.1 mV s⁻¹ (8-h cycle time) between the potential limits 0.3 V and −1.0 V (vs Hg/HgO) while recording the current drawn from the cathode at each voltage level.

Cathode performance was tested using galvanostatic cycling with potential limits (GCPL), in which the cathode was cycled under a constant current (C/3 rate) between two voltage limits (e.g., 0.3 V to −1.0 V vs Hg/HgO).

First, we tested whether layered Bi-birnessite could be electrochemically synthesized in situ in NaOH using the same recipe developed for synthesizing Bi-birnessite in KOH. We conducted electrochemical tests and materials characterization on EMD cathodes cycled once under four scenarios: (1) with Bi₂O₃ in NaOH; (2) with Bi₂O₃ in KOH; (3) without Bi₂O₃ in NaOH; and (4) without Bi₂O₃ in KOH. The final two scenarios were included to examine the role of Bi₂O₃ in the phase transition from EMD to birnessite.

We also compared the lifetime of the Bi-birnessite cathode, measured as total Ah processed per unit mass, in NaOH and KOH under the same use conditions. We conducted lifetime tests using the GCPL cycling protocol described above.

We note an inconsistency in the KOH concentration used for the experiment depicted in Fig. 5B. The concentration was 10.5 M KOH in Fig. 5B, rather than the 6.6 M used in the rest of the study. We have assessed the potential impact of this variation to be minimal. The pH increase between 6.6 M KOH (pH 14.8) and 10.5 M KOH (pH 15.0) is by approximately 1%, which is a relatively minor discrepancy. Thus, we assert that the difference in KOH concentration did not significantly affect the cathode lifetime outcomes reported.

Fresh and cycled materials were characterized using a P'Analytical X'Pert diffractometer with a Cu Kα radiation source. Samples were continuously scanned from 2θ = 10° to 80° with a step size of 0.026° and scan step time of 170 s. In cases where samples were pressed into stainless steel mesh, a monochromator was used to prevent fluorescence. Before XRD characterization, cycled electrodes were pressed under 4 metric tons cm⁻² to reduce low-angle scattering.

Scanning electron microscopy (SEM) was used to observe surface morphologies and phases present in samples. Samples were vacuum-dried after cycling and stored in ambient conditions in sealed plastic bags for one week prior to SEM analysis. SEM images were captured at 500 x, 10 kx, and 30 kx magnification using a Tescan Mira3 with a secondary electron detector at 30 kV.

We used X-ray fluorescence spectroscopy to determine the cation concentrations of cycled cathodes. The concentrations of Na and K cations were measured in cathodes that underwent one complete cycle and one full discharge after the formation of birnessite in NaOH and KOH, respectively. Cathodes that underwent a complete cycle were in the charged state and thus had the birnessite structure. Cathodes that only underwent one discharge were in the discharged state and thus had the Mn(OH)₂, or pyrochroite, structure.

Data were collected using an Oxford Instruments' XSupreme8000 with a He-filled test chamber. We calibrated the XRF data using two categories of standards: (1) EMD and potassium chloride and (2) EMD and sodium chloride. The calibration standards were prepared for each chloride using the weight percentages 0%, 0.5%, 1%, 1.5%, 2%, 5%, 10%, 20%, 30%, and 38%. Each calibration standard was ball-milled for 15 min. Calibration curves are shown in SI Fig. 6.

The process by which EMD is electrochemically converted to birnessite is illustrated in the voltage curves displayed in Figs. 1A–1D, and an annotated version is provided in SI Fig. 3. As reported by Yadav et al., during the first discharge of electrolytic MnO₂ (EMD) in KOH, MnO₂ converts into pyrochroite Mn(OH)₂ through a two-step reaction between manganese oxides and water. ^15,23

The first step is the solid-state protonation of EMD, which transforms into orthorhombic α-MnOOH, commonly referred to as groutite. ^1,14,15,24 This reaction can be observed as a sigmoidal shape in the discharge curve between 0.3 V and −0.4 V vs Hg/HgO. The second step is a dissolution-precipitation reaction, observed via the first discharge plateau at approximately −0.4 V vs Hg/HgO. ¹⁵

Upon subsequent charging, Mn(OH)₂ reverts to MnOOH at −0.2 V vs Hg/HgO, finally transforming to birnessite at 0 V vs Hg/HgO.

Our results show that the birnessite formation process was significantly influenced by the presence of Bi₂O₃ and the electrolyte constituents. When Bi₂O₃ was present (Figs. 1A and 1C), there were three significant plateaus during the first discharge and two plateaus during the subsequent charge. These five plateaus are presumably associated with the chemical reactions that transform EMD into layered. ¹⁵ Without Bi₂O₃ (Figs. 1B and 1D), the third discharge plateau at −0.6 V vs Hg/HgO and the first charge plateau at −0.5 V vs Hg/HgO vanish. It follows that these two plateaus correspond to Bi-redox reactions.

The voltage behavior at the end of the first charge is drastically different in the NaOH and KOH cases. In NaOH (Figs. 1A and 1B), there are two very small, successive plateaus at approximately −0.1 V and 0.1 V vs Hg/HgO. According to Yadav et al., these plateaus occur in the voltage range that corresponds to proton insertion in KOH systems. However, when the proton insertion reaction occurs in KOH, it results in a peak feature, as seen in the top right of Fig. 1C. The difference in the shape of the charge curves in NaOH and KOH indicate that different electrolytes generate different reaction mechanisms. Furthermore, the characteristic peak feature around 0 V vs Hg/HgO in KOH did not occur when Bi₂O₃ was absent. This suggests that the proton insertion reaction behavior in NaOH is unaffected by the presence of Bi₂O₃. In contrast, the proton insertion reaction in KOH is only detectable when Bi₂O₃ is present.

These findings merit future investigations of the impact of electrolyte properties (such as those listed in Table I) on birnessite formation and behavior.

Layered birnessites are intriguing as cathodes due to their 1 × ∞ tunnel structure, formed by sheets of edge-sharing MnO₆ octahedra. ²⁵ The space between the MnO₆ sheets often house water and various cations, which play a pivotal role in determining birnessite properties. A classic illustration is the larger spacing between K-birnessite layers than Na-birnessite, attributed to the size variations of the K⁺ and Na⁺ cations (see Table I). ²⁵

Structure: X-ray diffraction (XRD) was used to examine structural differences that arose during the birnessite formation cycle, based on the electrolyte composition and the presence of Bi₂O₃. Shown in Fig. 1E, the birnessite formation cycle resulted in prominent (001) peaks at 12°–13° and less intense (002) peaks at 25°–26°, regardless of which electrolyte was used and whether Bi₂O₃ was present. These peaks indicate that a layered birnessite structure formed in all of the tested conditions. However, the signal-to-noise ratios of the (001) and (002) peaks are slightly higher for the samples without Bi₂O₃. This suggests that the presence of Bi₂O₃ reduced the samples' birnessite concentration.

The samples' interlayer spacings, summarized in Table III, were noticeably impacted by the electrolytes and Bi₂O₃. Adding Bi₂O₃ increased the (100) interlayer spacing, d₁₀₀, by 1.1% in KOH and by 0.20% in NaOH. The additive also increased the (200) interlayer spacing, d₂₀₀, by 0.73% in KOH and by 0.15% in NaOH. Furthermore, switching from KOH to NaOH increased the interlayer spacings—d₁₀₀ expanded by 0.71% and 1.6%, with and without Bi₂O₃, while d₂₀₀ expanded by 0.47% and 1.1%, with and without Bi₂O₃. Given these trends, it is unsurprising that the smallest d-spacings occurred in KOH without Bi₂O₃ while the largest occurred in NaOH with Bi₂O₃. While it was expected that adding large Bi₂O₃ molecules would increase d-spacing, the expansion effect of the NaOH electrolyte was interesting since Na⁺ (2.04 Å in diameter) is smaller than K⁺ (2.76 Å in diameter). Early work by Post and Veblen similarly showed that the interlayer spacing in synthetic Na-birnessite was 2% higher than that in K-birnessite. ²⁶ Thus, the d-spacing is likely influenced by factors other than cation size, such as bonding characteristics like the solvation shell, which may also explain performance differences.

Table III. Comparative summary of interlayer spacings in samples cycled in KOH or NaOH, with or without Bi₂O₃.

Electrolyte	Active material	$2{\theta }_{100}$	${d}_{100}$ (Å)	$2{\theta }_{200}$	${d}_{200}$ (Å)
KOH	Bi-MnO2	12.6	7.04	25.1	3.51
NaOH	Bi-MnO2	12.5	7.09	25.0	3.56
KOH	MnO2	12.7	6.97	25.3	3.52
NaOH	MnO2	12.5	7.08	25.0	3.55

Composition and Electrolytic cation incorporation: The ratio of intercalated electrolytic cations to MnO₂ was analyzed using X-ray fluorescence spectroscopy (XRF). As summarized in Fig. 1F, the molar concentration of Na in Na_xMnO₂ in the charged and discharged states was 0.19 and 0.025, respectively. The molar concentration of K in K_xMnO₂ in the charged and discharged states was 0.23 and 0.00, respectively. The increase in cation concentration in the charged cathodes indicates that Na and K both inserted into the layered birnessite structure during charging. The molar fraction of K was 24% higher than that of Na, suggesting K inserts into the cathode more effectively.

While little is known about the impact of cation content on the electrochemical performance of birnessite, recent work by Gao et al. studied the impact of cation content on nanostructured pseudocapacitive birnessite electrodes. ²⁷ The authors found that an increase in Mn vacancy concentrations, which was associated with higher cation incorporation, served as low-energy intercalation sites and thereby increased the specific capacitance of the electrode. Thus, the performance trend observed by Gao et al.—in which higher cation concentration corresponded with improved performance–was opposite to the trend we observed in this work—in which lower cation concentration correlated with improved performance. However, future research could adapt the authors' approach and insights to more deeply explore the impacts of ion incorporation on birnessite performance.

Morphology: SEM was used to compare the influence of the KOH and NaOH electrolytes on the morphological evolution of the cathode during cycling (SI Fig. 8). For comparison, the morphology of a fresh, uncycled cathode is shown in SI Fig. 7. During the initial birnessite formation cycle, significant morphological differences were evident. After the first discharge, low magnification images showed that bright clusters were distributed throughout the KOH sample (SI Fig. 8A) but were absent from the NaOH sample (SI Fig. 8B). The bright clusters were flower-like agglomerations of 1–2 μm flake- and needle-like particles. While similar flower-like clusters have been associated with birnessite in previous work, ^28,29 they were not found to have been reported for the discharged, pyrochroite phase during MnO₂ cycling. Instead, Patrice et al. observed cubic particles after discharging a 3 wt% γ-MnO2 cathode in KOH to form pyrochroite. ³⁰ Similar orthorhombic-like particles were observed in this work—2–3 μm wide in KOH and 0.6–1.2 μm wide in NaOH—dispersed within a dense matrix of significantly smaller (~0.2 μm) particulate features. It is worth considering that the SEM images after discharge may partially reflect changes due to the air exposure prior to imaging, which could lead to Mn(OH)2 oxidation.

After the first charge, which completes the birnessite formation cycle, the birnessite in the KOH sample (SI Fig. 8C) resembled previously observed assembled birnessite nanosheets. ^2,27,31 On the other hand, the birnessite in the NaOH sample predominantly displayed small 0.5–1.0 μm bulk particles (SI Fig. 8D) and occasional large hexagonal flakes measuring 1–2 μm (seen in the figure insets). Large 2 by 5 μm particles, likely the tetragonal form of Bi₂O₃ detected via XRD, were also evident in the NaOH sample.

By the second cycle, similarities in cathode morphologies emerged between the two electrolyte environments. After the second discharge, the cathodes in both electrolytes presented predominantly smooth morphologies interspersed with 200 nm bulk particle clusters and sporadic larger 5 μm particles electrolytes (SI Figs. 8E–8F). The subsequent charged sample in KOH was characterized primarily by 200 nm bulk particles (SI Fig. 8G), while the NaOH sample displayed a mix of smooth portions and occasional 5 μm needle-like particles (SI Fig. 8H).

Given the interesting morphological differences observed in this work, future work should thoroughly study relationships between the electrolyte environment and the cathode's morphological evolution and performance.

We assessed the cathode's stability at full depth-of-discharge (DOD) by cycling it between 0.3 V to −1.0 V vs Hg/HgO at a C/3 rate. The lifetime test concluded when the cathode's capacity declined by 80%. We defined cathode lifetime as the total Ah processed per unit of active mass. Additionally, we examined cathode performance through cyclic voltammetry, with a voltage sweep rate of 0.1 mv s⁻¹ within voltage limits of 0.3 V to −1.0 V vs Hg/HgO.

The discharge and CV curves in Fig. 2 showcase the cathode's behavior during its first 15 cycles in a flooded beaker cell using NaOH (A)–(B) and KOH (C)–(D) electrolytes. Fast capacity fade occurred in both KOH and NaOH. After 15 cycles, the discharge curves (Figs. 2A and 2C) show that NaOH maintains marginally superior capacity; yet the CV scans (Figs. 2B and 2D), while very similar, show that KOH has a 18% higher discharge peak and 1.2% larger area under the curve. This contrast might arise from differing reaction kinetics between the electrolytes. While KOH seems better for quick reaction, evident in CV scans, NaOH appears more stable during the slower reactions of cycling. Together, the CV and cycling plots provide a nuanced view of the electrolytes' electrochemical differences.

Upon closer examination, the discharge plateau at −0.4 V in Figs. 2A and 2C stands out as a distinctive comparison between the two electrolytes. Corresponding to peak (i) in Figs. 2B and 2D, this plateau signifies the dissolution-precipitation reaction converting MnOOH to Mn(OH)₂. While this reaction is short-lived in KOH, disappearing after 4 cycles, it persists in NaOH for at least 10 cycles. Thus, NaOH appears to sustain the reaction more sustainably.

The −0.6 V discharge plateau and its corresponding peak (ii) also draw attention as they represent the Bi-redox processes. Notably, this reaction exhibited improved stability in both electrolytes, compared to the other reaction steps. Figure 2A reveals that this plateau is noticably more constant in NaOH than KOH.

Shifting focus to another key comparison, the behavior of peak (v) at approximately 0.15 V is drastically different between the two electrolytes. During the first few cycles, peak (v) is very pronounced in KOH CV curve (Fig. 2D), corresponding to a short charge plateau (Fig. 2C) associated with solid-state protonation of MnOOH. In contrast, this peak is barely detectable in the NaOH CV scan (Fig. 2B). Evidence of this reaction occurring in NaOH is more prominent in the charge curves (Fig. 2A), which shows two small successive plateaus in the V-range of interest. This contrast in peak behavior suggests that birnessite formation occurs via slower kinetics in NaOH than in KOH.

Finally, the rightmost bar in Fig. 3A presents the cathode lifetimes at full DOD in KOH (blue) and NaOH (yellow). The NaOH system had a lifetime of 13.7 Ah g⁻¹, outperforming KOH in the beaker cell system with a 70% longer cathode lifetime.

To further our understanding of the system's performance in each electrolyte, we analyzed the influence of the lower voltage limit during cycling on the cathode's lifetime. While the upper voltage limit was set to 0.3 V vs Hg/HgO, we varied the minimum voltage to which the cathode was discharged. In each test, the upper and lower voltage limits remained fixed throughout the lifetime of the cathode. The lifetime test concluded when the cathode's capacity fell below 20% of its initial capacity at the predefined voltage range. Cathode lifetime was calculated as the total Ah processed per gram of MnO₂.

Figure 3A visualizes the relationship between cathode lifetime (in Ah throughput) and lower voltage limit in both electrolytes. Similarly, SI Fig. 10 shows how cycle life is influenced by the lower voltage limit. Each bar in the figure represents a different cathode lifetime test. Our analysis revealed a weak positive correlation between lifetime (in Ah/g) and lower voltage limits (i.e., deeper discharge). The correlation was stronger in KOH (R² = 0.67, p-value = 0.024) than in NaOH (R² = 0.55, p-value = 0.055). To summarize this trend, we calculated the percentage increase in lifetime for each electrolyte, M (where M represents K or Na), using the formula: % lifetime increase in MOH = $({\rm{x}}{\rm{-}}{\rm{y}})/{\rm{x}},$ where ${\rm{x}}$ is the lifetime in MOH when the minimum voltage is −1.0 V, and ${\rm{y}}$ is the lifetime when the minimum voltage is −0.4 V. Remarkably, this approach revealed that lifetimes surged by 175% in NaOH and 487% in KOH when discharged to the deepest level (−1.0 V vs Hg/HgO), compared to the highest minimum voltage tested (−0.4 V vs Hg/HgO).

In every voltage window tested, NaOH outperformed KOH. This superiority is especially clear in Fig. 3B, which illustrates the percentage increase in cathode lifetime when cycled in NaOH relative to KOH. The figure indicates a 20% to 250% lifetime enhancement in NaOH. The most pronounced divergence in lifetimes (over 200%) occurred when the minimum voltage was −0.5 or −0.6 V vs Hg/HgO. Figures 4A and 4B show the immediate capacity retention improvement in NaOH from the beginning of the tests. Figure 4C further highlights the significantly higher 10th discharge capacity in NaOH. Importantly, as showcased by Fig. 4C, setting the minimum voltage to −0.5 and −0.6 V vs Hg/HgO omits the Bi-redox plateau. This suggests that lifetime boost in NaOH mainly stems from the steadier dissolution-precipitation reaction. The dissolution-precipitation plateau decayed immediately in KOH when the cathode was discharged −0.5 or −0.6 V, disappearing after just 3–4 cycles (as seen in the 2nd and 3rd columns in Fig. 4C). In contrast, the same reaction plateau showed significantly higher stability in NaOH.

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