Electrolyte Optimization of a Dual Compartment Hydrogen Peroxide Fuel Cell with Prussian Blue and Tantalum Electrodes

Abstract

Hydrogen peroxide fuel cells have emerged as a promising class of electrochemical energy conversion device owing to the dual redox character of H₂O₂, its liquid-phase storage, and its ability to operate in air-free environments. In this work, a dual-compartment direct H₂O₂ fuel cell using a Prussian Blue cathode and a tantalum anode, separated by a Nafion 115 proton exchange membrane, was systematically characterized and optimized with respect to electrolyte pH and ionic composition. The influence of pH on OCV was investigated independently in each compartment across the range of pH 2 to 12. In the tantalum compartment, OCV increased non-linearly with pH from 573 mV to 808 mV, driven by the enhanced electrochemical reactivity of the system under alkaline conditions. In the Prussian Blue compartment, OCV decreased from 676 mV to 199 mV with increasing pH, reflecting the instability of the material in alkaline conditions. The effect of the electrolyte ionic composition on average current density was subsequently investigated by varying the concentrations of NaCl and Dy(NO₃)₃. Increasing NaCl from 0 to 2.5 M produced an increase in current density from 0.414 mA/cm² to 0.973 mA/cm², consistent with ohmic resistance reduction through improved ionic conductivity. The addition of Dy(NO₃)₃ produced a positive response with an optimal concentration of $0.05$ M, at which current density reached 1.08 mA/cm², before declining sharply. Under the fully optimized conditions, pH 12 in the tantalum compartment, pH 2 in the Prussian Blue compartment, 0.3 M H₂O₂, 2.0 M NaCl, and 0.05 M Dy(NO₃)₃, the cell produced an OCV of 724 mV and a peak power density of 0.283 mW/cm² at a current density of 0.8 mA/cm². These results demonstrate that meaningful electrochemical performance can be achieved in a dual-compartment H₂O₂ fuel cell without the use of precious metal catalysts and highlight electrolyte engineering as an effective strategy for improving cell output in this class of device.

1. Introduction

The global transition away from fossil fuels has increased interest in researching alternative energy sources like electrochemical technologies. Fuel cells convert chemical energy directly into electrical energy through oxidation and reduction reactions, offering high theoretical efficiencies with almost zero emissions. Conventional hydrogen–oxygen (H₂-O₂) proton-exchange membrane fuel cells (PEMFCs)are the most widely studied and established systems. In these cells, H₂ gas is oxidized at the anode and O₂ is reduced at the cathode, yielding water as the only byproduct. Despite the high efficiencies and near-zero emissions, PEMFCs face significant practical barriers related to storage and transportation. Other established fuel cell technologies present their own distinct challenges. For instance, solid oxide fuel cells (SOFCs) require extremely high operating temperatures, and phosphoric acid fuel cells (PAFCs) rely on highly corrosive concentrated acids. Even alternative systems that utilize liquid fuels, such as Direct Methanol Fuel Cells (DMFCs), face severe drawbacks like methanol crossover through the membrane, a reliance on toxic fuel, and the emission of carbon dioxide (CO₂) as a byproduct [1].

In contrast, hydrogen peroxide (H₂O₂) fuel cells have gained interest as a promising alternative that circumvents many of these operational hurdles, clearly outlining the potential advantages of investigating this chemistry. Unlike gaseous H₂, H₂O₂ is stable as a liquid and soluble in water, making it significantly easier and safer to store and transport without the need for high-pressure infrastructure. Furthermore, unlike DMFCs, H₂O₂ systems offer a truly zero-emission profile, producing only water and oxygen. H₂O₂ offers a higher standard reduction potential than oxygen (1.776 V vs. SHE for H₂O₂ and 1.229 V vs. SHE for O₂ in acidic medium). Because it also has the ability to act as the oxidant and the reductant, it can remove the need for a traditional ion-exchange membrane. This drastically reduces the complexity and cost of the fuel cell compared to SOFC or PAFC systems, while completely bypassing the crossover issues that plague DMFCs. The reactions at the anode and cathode are as follows:

$$\mathrm{Anode}:{\mathrm{H}}_2{\mathrm{O}}_2\to {\mathrm{O}}_2+2{\mathrm{H}}^{+}+2{\mathrm{e}}^{-}$$

(1)

$$\mathrm{Cathode}:{\mathrm{H}}_2{\mathrm{O}}_2+2{\mathrm{H}}^{+}+2{\mathrm{e}}^{-}\to 2{\mathrm{H}}_2\mathrm{O}$$

(2)

$$\mathrm{Total}:2{\mathrm{H}}_2{\mathrm{O}}_2\to {\mathrm{O}}_2+2{\mathrm{H}}_2\mathrm{O}$$

(3)

In a single-compartment configuration, both reactions occur in the same electrolyte and, depending on the combination, will have catalytic activity in a specific direction. In a dual-compartment configuration, each compartment is separated by a proton exchange membrane, where the anodic and cathodic electrolytes are maintained independently. Therefore, the parameters of the electrolyte, like the pH and composition of each compartment, can be optimized for its electrode. There have been a variety of experiments researching H₂O₂ fuel cells. They can focus on testing new electrodes, membranes, and electrolyte compositions, as seen in the

Table 1. Summary of recent experimental studies on H₂O₂ fuel cells, detailing cell configuration, electrode materials, electrolyte composition, and performance metrics.

Year	Cell Type	Anode	Cathode	Electrolyte	OCV (mV)	PPD (mW/cm2)	References
2024	Dual	Graphite Felt	Pt/C	e-fuel/H2O2	1180	72.9	[2]
2024	Single	BP/PEI /CoPc	BP/PEI /hemin	H2O2	311	0.232	[3]
2022	Single	Graphite Felt	Pt/C	e-fuel/H2O2	1120	1456	[4]
2022	Dual	Ni	PB/CNT	H2O2(alkaline)/ H2O2(acidic)	1000	8.9	[5]
2016	Dual	Unspecified	PB/ CNT	Methonol /H2O2	740	125	[6]

.

The electrode combination of Prussian Blue and Tantalum (Ta), paired with a dysprosium salt additive, was identified by Alderson [7]. The author tested various pure metals, metal–organic complexes, and salt additives to identify the strongest performing combinations in a single compartment configuration. This paper aims to expand on that research by testing this electrode combination in a dual-compartment configuration and optimizing the electrolyte in each compartment for each electrode. This is done by measuring the influence that the electrolyte pH has on the open circuit voltage (OCV) of the cell. Then, to optimize the current density, the support electrolyte composition is tested to see what concentration produces the best performance in the cell.

2. Materials and Methods

2.1. Materials and Instrumentation

The chemical materials used throughout the experiments are listed in

Table 2. Materials used for the experiments with their source and purity.

Compound/Material	Source	Purity	CAS
Dysprosium nitrate hydrate (Dy(NO3)3)	Sigma-Aldrich (Burlington, VT, USA)	99.9%	100641-13-2
Hydrochloric acid (HCl)	Sigma-Aldrich (Burlington, VT, USA)	37%	7647-01-0
Hydrogen peroxide (industrial grade) (H2O2)	Arkema (Columbes, France)	50%	7722-84-1
Iron (III) ferrocyanine (Prussian blue)	Millipore Sigma (Burlington, VT, USA)	Not disclosed	14038-43-8
Nafion N115 Membrane	Ion-Power (Tyrone, PA, USA)	N/A	31175-20-9
NanoSilvex P-01 conductive silver paste	NanoCnet (Kitchener, Ontario, CA , USA)	N/A	N/A
Polishing paste, 5000 grit (2–3 µm)	TechDiamondTools (Santa Clara, CA, USA)	N/A	N/A
Polishing paste, 14,000 grit (1–0.5 µm)	TechDiamondTools (Santa Clara, CA, USA)	N/A	N/A
Polishing paste, 50,000 grit (0.5–0.01 µm)	TechDiamondTools (Santa Clara, CA, USA)	N/A	N/A
Sodium chloride (NaCl)	Bioshop Canada Inc. (Burlington, CA, USA)	99%	7647-14-5
Sodium Hydroxide (NaOH)	Bioshop Canada Inc. (Burlington, CA, USA)	98%	1310-73-2
Tantalum (foil)	Alfa Aesar (Ward Hill, MA, USA)	99.95%	7440-25-7

. Standard laboratory glassware and equipment, including beakers, volumetric pipettes, a scale, and a magnetic stirrer, were used. Cell voltage was recorded using a Titan S8 Data Logger (MadgeTech, Warner, NH, USA), and solution pH was measured with an Apera Instruments AI-312 pH meter (Apera Instruments, Columbus, OH, USA). The full specifications for both instruments are provided in

Table 3. The specifications of measurement devices used in the experiments.

Device	Measurement	Resolution	Accuracy
Titan S8 Data logger	Voltage	0.001 mV	±0.1 mV
Apera Instruments LLC-AI312 pH meter	pH	0.01	$\pm 0.01$

. Prior to use, the pH meter was calibrated using a two-point calibration with pH 4 and pH 7 buffer solutions.

2.2. Electrode Fabrication

The Prussian Blue electrode material was prepared by dissolving 38.6 mg of the compound in 100 $\mathsf{\mu }$L of a base solvent compared of 700 parts Dimethylformamide (DMF), 100 parts Tetrahydrofurane (THF), 200 parts 70% isopropanol, and 0.1 mg lampblack carbon in a glass vial. The solution was applied to carbon fiber cloth and dried in an oven over two applications.

The electrodes were then fabricated by cutting the electrode material into a 2 cm by 3 cm rectangle. Copper tape was applied at the contact region to increase the effective joint surface area, and silver paste was applied between the copper tape and the electrode material to minimize contact resistance and improve electrical conductivity. A length of copper wire was then attached, with a spade connector soldered to one end to provide a secure and repeatable connection to the data logger terminals. The continuity and resistance of each completed electrode assembly were verified using a digital multimeter. There was a marginal increase in resistance relative to the electrode material, which is expected and deemed acceptable for the purposes of these experiments.

2.3. Cell Configuration

All pH and concentration experiments were conducted using an H-cell membrane reactor with a Nafion 115 proton exchange membrane separating the two electrolyte compartments. The H-cell is pictured in Figure 1.

2.4. Optimizing pH Procedure

The influence of electrolyte pH on cell OCV was investigated across the range of pH 2 to 12. The electrolyte in each compartment consisted of 0.3 M H₂O₂, 1.5 M NaCl, and 0.05 M Dy(NO₃)₃. Both compartments were initially prepared at pH 2. The pH of one of the compartments was then adjusted incrementally using acidic and alkaline solutions, prepared by dissolving 5 g of HCl or NaOH, respectively, in 40 g of deionized water, while keeping the other compartment at a pH of 2. At each pH increment, the cell was allowed to stabilize before the output voltage was recorded over a 5 min measurement window. Once all the measurements were complete, the first compartment was held constant and the second compartment was adjusted. This experiment was repeated three times for reproducibility.

2.5. Optimizing Electrolyte Composition Procedure

The effect of electrolyte composition on cell current density was evaluated in two sequential steps. In the first step, the NaCl concentration was varied from 0 M to 2.5 M in increments of 0.5 M to identify the optimal concentration of NaCl. In the second stage, having established the optimal NaCl concentration, Dy(NO₃)₃ was added at concentrations ranging from 0 M to 0.1 M in increments of 0.025 M to assess its additional contribution to cell performance. Each concentration was evaluated three times for reproducibility. For all concentration experiments, the electrolyte in both compartments was freshly prepared for each measurement and contained a base of 0.3 M H₂O₂ and 0.12 M HCl. The NaCl and Dy(NO₃)₃ concentrations were then adjusted as described above.

In both the pH optimization and the electrolyte composition experiments, a stabilization period of about 5 min was used prior to data recording, followed by the measurement collection over a 300–350 s window. The stabilization period allowed for the electrode and electrolyte to reach a quasi-steady state condition before the measurements were recorded. Data collection only commenced once a visual plateau in the output (potential or current) was observed in the data logger display. This visual check provided an additional metric to ensure measurements were not recorded during periods of active drift.

The recording window was selected on the basis that the OCV and current density signals had plateaued and did not have a noticeable drift within this period under the majority of the conditions tested. This approach was adopted primarily because the study contained a large number of experimental conditions across two experimental stages. There were eleven levels of pH variation, with each level tested three times, and six NaCl and eight Dy(NO₃)₃ concentration variations were also tested three times. An extended stabilization and recording period was impractical within this experimental framework. A prolonged operation may introduce additional degradation mechanisms, including H₂O₂ depletion, which would cause the reported values to deviate from the steady state performance.

2.6. Polarization Curve Procedure

The polarization curves were acquired using a galvanostatic method with a DC load generator. The H-cell membrane reactor was used with the optimized pH and electrolyte composition. The current density was increased incrementally with the load set to draw at a fixed current at each measurement point. The cell voltage was measured and recorded for a period of 3 min and then averaged. The power density at each point was calculated as the product of the measured cell voltage and the applied current density.

3. Results and Discussion

3.1. The Optimal Compartment pH Results

First, looking at the influence of pH on the Ta compartment, Figure 2 shows that the cell voltage increased non-linearly with pH, rising from 573.7 mV at pH 2 to 808.3 mV at pH 12, a net increase of 40.9%. The relationship was not strictly linear, a local minimum was observed at pH 6, followed by a recovery and sustained increase from pH 7 to 12. Figure 3 shows that the cell voltage remained stable across most electrolyte pH conditions over the 300 s measurement period. The notable exceptions include pH 6, where the cell voltage fluctuated persistently, and pH 5 and 8, which both had drifted upwards before stabilizing. The remaining pH conditions reached and maintained a stable voltage for the duration of the measurement.

The increase in OCV with pH is consistent with the Nernst equation when applied to the half-reactions governing the electrodes. Jerkiewicz explained for the H⁺(aq)/H₂(g) redox in acidic media, the application of the Nernst equation shows that the electrode potential decreases by 59.2 mV per unit increase in pH, since $pH=-loga\left({\mathrm{H}}^{+}\right)$ and the proton activity directly enters the expression for equilibrium potential. An analogous relationship exists in alkaline media, where the redox couple becomes H₂O(aq)/H₂(g),OH⁻(aq). The corresponding Nernst expression’s dependence on hydroxide ion activity can be expressed as shown in Equation (4). The application of this equation yields the same plot as with the acidic media [8]. As the half-cell potential reaction at the Ta anode site decreases with pH, and the overall cell potential is the difference between the potential of the cathode site and anode site, there would be an overall increase in cell potential. A similar potential decrease with pH was observed by Macagno and Schultze when quantifying the growth of the oxide produced by Ta, Tantalum pentaoxide (Ta₂O₅). It was observed that pH shifted 60 mV per pH unit [9].

$$E=E^{\circ }-{\displaystyle \frac{RT}{F}}ln(\sqrt{f\left({\mathrm{H}}_2\right)/p^{\circ }}\times a\left({\mathrm{O}\mathrm{H}}^{-}\right))$$

(4)

The observed drop in cell potential at pH 6 is a deviation from the upward-trending relationship between pH and OCV. This occurrence was reproduced across 3 independent measurements, confirming it was not due to experimental error. A similar dependence of electrochemical behavior on pH was reported by Attarzadah et al. [10]. They observed a non-linear variation in the corrosion current density and passive film resistance of Ta in a borate buffer solution across a comparable pH range. The optimal protective behavior was at pH 8.0. This illustrates that non-linear pH responses are not without precedent in electrochemical systems. The precise reason for the deviation at pH 6 in the experimental configuration cannot be determined by electrochemical measurements alone. Future investigations using surface characterization techniques like X-ray photoelectron spectroscopy (XPS), scanning electron microscopes or energy-dispersive X-ray spectroscopy (EDS) would be needed to understand the underlying mechanisms.

H₂O₂ is well known for its ability to undergo a disproportionation reaction, where the peroxide breaks down into water and O₂ spontaneously. This reaction provides the electrolyte a pathway to react without being electrochemically productive as the H₂O₂ is spent without reacting with the electrode material. This reaction is known to be accelerated under alkaline conditions as well as on the metal oxide surfaces. As both have high pH, and a tantalum oxide passive film is present in the experiment, disproportionation may interfere with the performance results of the cell.

During the experiments, bubble production was observed through the pH levels. There was no notable increase in bubble production at the higher pH values; without a direct measurement of the H₂O₂ consumption rate, it is not possible to determine what fraction of the H₂O₂ was consumed productively through the electrochemical reactions versus unproductively through chemical disproportionation. The output represents a net output of the cell from the combined influence of both pathways. This leaves an opportunity to conduct future experiments to find the faradic efficiency, which can be completed using periodic sampling and titration of the electrolyte or an assay, and comparing the consumption rate to the integrated current. This would provide a basis for assessing the productive utilization of H₂O₂ by the cell.

After the Ta runs, the Prussian Blue compartment’s pH was adjusted from pH 2 to 12. The OCV of the cell was measured and is shown in Figure 4. The cell voltage decreased approximately linearly with increasing pH, falling from 676.4 mV at pH 2 to 199.7 mV at pH 12, a net decrease of 70.5%. Unlike the tantalum compartment results, the relationship was consistent. There were no significant variations observed across the pH range. Figure 5 shows that the cell voltage remained relatively stable across all pH conditions over the 300 s measurement period, though a gradual downward drift was observed across most pH levels.

The Prussian Blue compartment’s OCV response was the opposite of the Tantalum compartment. This behavior is most likely due to Prussian Blue’s instability in alkaline conditions. Karayakin et al. demonstrated that Prussian Blue modified electrodes were stable in neutral and acidic supporting electrolytes, but the material degraded progressively in alkaline solutions [11]. The dissolution of Prussian Blue in the alkaline media would result in the loss of material from the electrode surface, resulting in a loss of the reduction reaction sites.

It is important to note that there are two distinct sources of potential that contribute to the measured OCV of the pH-optimized cell. First is the electrochemical fuel cell potential created from the oxidation and reduction of H₂O₂ at the anode and cathode. There is a secondary cell potential that is generated by the junction potential caused by the proton activity gradient from the varying pH levels on either side of the membrane. The theoretical contribution of a concentration cell with a pH difference of 10 units is about 0.592 V at 298.15 K, as given by the Nernst equation when applied to a pH gradient. The total measured OCV was 730 mV in the optimized conditions, and it began with a baseline of 522 mV when both compartments were at pH 2. In an ideal sense, if the entire contribution of the concentration cell was added to the baseline, the cell would have an overall potential much larger than 1 V. As the measured OCV was below that, it is not solely the contribution from the increase in potential. Further investigations will be required comparing the cell at its optimized pH levels with and without H₂O₂ to identify the specific contribution of the concentration cell.

Additionally, a large pH gradient between the membrane may drive a proton crossover from the acidic to the alkaline compartment, progressively neutralizing the pH gradient. This self-consumption of the acid and base would be independent of the H₂O₂ consumption and could be a source of performance degradation under sustained operation. Since the experiments were measured over a 300–350 s window, the degree of collapse of the pH gradient is considered to be small, and the reported values are a reasonable representation of the initial pH gradient. For future sustained operation, strategies like continuous electrolyte replenishment can be employed to minimize the effect of the acid–base self-consumption.

Therefore, from the standpoint of maximizing the cell potential, the results indicate that the configuration with the Tantalum compartment at a pH value of 12 and the Prussian Blue compartment at a pH of 2 should yield the highest achievable cell potential.

During the pH experiments, the addition of NaOH to the electrolyte produced a white precipitate, as shown in Figure 6. The quantity of precipitate increased with further NaOH addition. Based on the chemical reactions occurring in the system, this precipitate is believed to be a combination of dysprosium hydroxide (Dy(OH)₃) and NaCl. Dysprosium is a lanthanide known for its ability to be dissolved in aqueous solutions and precipitate upon the addition of aqueous alkali [12]. The dissolved Dy³⁺ reacts with the NaOH to yield solid Dy(OH)₃, while the NaOH simultaneously reacts with the HCl present in the electrolyte to produce NaCl in solution, as described by the reactions in Equations (5) and (6). The addition of HCl and NaOH to the electrolyte and the formation of the precipitate changes the electrolyte composition, which introduces some additional influences on results. The HCl and NaOH concentrations were chosen to balance pH control without diluting the electrolyte solution. Future work should include attempting the experiment with other acidic and alkaline solutions to minimize the effect of other variables.

$$\mathrm{Dy}{\left({\mathrm{NO}}_3\right)}_3+3\mathrm{NaOH}\to \mathrm{Dy}{\left(\mathrm{OH}\right)}_3+3{\mathrm{NaNO}}_3$$

(5)

$$\mathrm{HCl}+\mathrm{NaOH}\to \mathrm{NaCl}+{\mathrm{H}}_2\mathrm{O}$$

(6)

3.2. The Optimal Electrolyte Composition Results

The cell current density was measured as a function of NaCl concentration in the electrolyte across the range of 0 M to 2.5 M. As shown in Figure 7, the cell current density increased with NaCl concentration, increasing from 0.414 mA/cm² at 0 M to 0.973 mA/cm² at 2.5 M, an overall increase of 136.6%. The relationship is approximately linear, though a plateau was observed between 0.5 M and 1 M, followed by an increase from 1.5 M onward before leveling off again at 2.5 M. Figure 8 shows the cell current density over the 300-s measurement period at each concentration. The 0 M and 1.5 M conditions had a gradual downward trend throughout the measurement, while the other concentrations, particularly 2 M and 2.5 M, remained comparatively stable.

Figure 7 and Figure 8 clearly present a positive relationship between the NaCl concentration and the cell current density. This relationship continues up to about 2 M, after which the current begins to plateau. This behavior is expected by the supporting electrolyte in electrochemical cells. A supporting electrolyte in a solvent is used to support the electron transfer between the electrode compartments. A large supporting electrolyte concentration will increase the solution conductivity by reducing the solution resistance [13]. This is seen in experiments: as the NaCl concentrations increase, so does the current density.

Schalenbach et al. examined the impact of supporting salts on ion transport and limited currents in electrolyte and ionic liquids and established a physical model that described how increasing the salt concentration reduces the transport of electroactive species while also lowering the overall resistance of the cell [14]. It presents a compromise between the increased conductivity and the limited current of the cell. The experiments in the H₂O₂ fuel cells map to this conclusion. At low NaCl, the cell has a higher internal resistance, and the addition of NaCl improves the performance of the cell by reducing that resistance. The plateau observed past 2 M of NaCl is consistent with this model, as at high salt concentrations, the shared conductance between the ions lowers transport of the electrochemically active ions, limiting further increases in the current density. Therefore, to avoid the plateau, the optimal NaCl concentration is 2 M. The observed increase in current density with NaCl concentration is understood to be primarily due to a reduction in bulk solution resistance through improved ionic conductivity. This is consistent with the established role of supporting electrolyte in electrochemical cells as described by Schalenbach et al. However, it is important to note that the total ohmic resistance of a cell consists of multiple components beyond bulk electrolyte resistance, including the resistance of the proton exchange membrane, the electrode-electrolyte interface, and resistance in the electrode assembly. To further investigate the resistance components, a method like electrochemical impedance spectroscopy (EIS) will need to be employed in future work.

The chemical compatibility of the optimized electrolyte with the electrode material is an important consideration for the sustained operation of the cell. Hypochlorous acid can form through the oxidation of Cl⁻ by H₂O₂, which can contribute to the degradation of the Prussian Blue electrode. However, due to ambient temperature, the absence of applied potentials in excess of the OCV, and the relatively dilute H₂O₂ concentration, the rate of Cl⁻ oxidation may be limited. The Cl⁻ may also promote the localized breakdown of the Ta₂O₅ film, but a reproduceable trend in the tantalum component suggests that the film remained intact throughout the experiment duration. The competitive adsorption of Cl⁻ can inhibit the H₂O₂ reduction activity at the Prussian Blue surface. As the current density increased with concentration until 2 M of NaCl, this effect may be more pronounced above this concentration. Post-experiment surface characterization using techniques like SEM, EDS and XPS will be employed in future experiments to assess and evaluate the chemical compatibility of the electrolyte and electrode material.

The cell current density was measured as a function of Dy(NO₃)₃ concentration in the electrolyte across the range of 0 M to 0.1 M, with the NaCl concentration fixed at 2 M. As shown in Figure 9, the cell current density initially increased with Dy(NO₃)₃ concentration, reaching a peak of approximately 1.08 mA/cm² at 0.05 M, before declining at higher concentrations to 0.68 mA/cm² at 0.1 M. Figure 10 shows the cell current density over the 300 s measurement period at each concentration. The 0.05 M and 0.0375 M conditions produced the highest current densities and remained comparatively stable, while the 0.0625 M, 0.075 M, and 0.1 M conditions were notably lower and less stable.

Figure 9 and Figure 10 show that with the addition of Dy(NO₃)₃, there is a modest performance increase from 0.951 mA/cm² at 0 M to a peak of 1.08 mA/cm² at 0.05 M, followed by a sharp decline to 0.68 mA/cm² at 0.1 M. Dy(NO₃)₃ has been previously identified to be a strong additive to increase to the OCV of the cell [7]. Beyond its contribution to increase the rate of reduction, as with NaCl, the addition of Dy³⁺ cation may contribute to electrochemical conductivity through its interaction with the electrode surface. Briega-Martos et al. discusses how cations in solution can tune the structure of the interfacial water adlayer at the electrode surface, influencing the rate of the H₂O₂ reduction reaction on Pt(111) electrodes [15]. Although the authors focused on alkali metal cations, the same mechanism is expected to apply to lanthanides like Dy³⁺, considering they are a higher charged ion. Furthermore, lanthanides such as Dy are known for their unique 4f electronic structures, whose movement is known for its ability to decrease activation energy and improve catalytic processes [16].

Figure 11 presents the polarization and power density curves of the dual-compartment H₂O₂ fuel cell using the optimized electrolyte in the H-cell membrane reactor. The cell exhibited an OCV of 724 mV, consistent with the results in pH experiments. As the current density increased, the cell potential decreased steadily, following a typical polarization behavior expected of an electrochemical fuel cell. The power density increased with current density, reaching a peak of approximately 0.283 mW/cm² at a current density of 0.8 mA/cm². After the peak, a sharp decline is observed which is characteristic of concentration polarization, where mass transport limitations restrict the supply of H₂O₂ to the electrode surface. The cell reached a limiting current density of approximately 1.1 mA/cm², at which point the cell potential approached zero. As a comparison, Figure 12 shows the polarization curve of the H-cell using an electrolyte of 0.3 M H₂O₂, and 0.5 M NaCl, with the Ta compartment and the Prussian Blue compartment both at a pH of 2. The OCV if the cell was 522 mV, and it had a peak power density of about 0.101 mW/cm² at 0.5 mA/cm².

The polarization and power density curves were generated from the H-cell configuration using the optimized electrolyte consisting of 0.3 M H₂O₂, 0.05 M Dy(NO₃)₃, 2.0 M NaCl, and Ta and Prussian Blue compartments at pH 12 and 2, respectively. Figure 11 depicts the characteristic three-region profile expected from an electrochemical fuel cell [17].

The initial voltage drop from the OCV of 724 mV as the current density increased is attributed to the activation overpotential of the cell. This is the kinetic barrier that the cell has to overcome to drive the H₂O₂ reactions in the Ta and Prussian Blue compartments. The activation overpotential is described by the Butler–Volmer equation and is characterized by an exponential drop in cell voltage in the low current densities [18]. At the intermediate current densities, the polarization curve transitions to an ohmic loss region where the voltage decreases with increasing current density. The sources of ohmic resistance in this cell would include the PEM membrane and resistances in the electrode [19]. The mass transport region occurs at about 0.80 mA/cm². This is also where the peak power density of about 0.283 mW/cm² was achieved. This occurs when the H₂O₂ is consumed at the electrode, and the rate at which it is consumed begins to exceed the rate at which it can be replenished by the surrounding electrolyte. This current density was limited to 1.1 mA/cm², meaning that beyond that load, no further electrochemical reaction can occur.

Figure 12 is the result from the H-cell without any of the optimizations made to the electrolyte. The cell had an OCV of about 522 mV with a peak power density of 0.101 mW/cm² at 0.5 mA/cm². The optimization of the electrolyte had a direct benefit to the output, as both the potential and current density increased due to the changes.

4. Conclusions

This study presents an experimental investigation of the electrolyte conditions that govern the performance of a dual-compartment H₂O₂ fuel cell using a Ta anode with a Prussian Blue cathode.

The OCV of the cell was heavily dependent on the electrolyte pH of each compartment. The tantalum compartment OCV increased with the pH from 573 mV at pH 2 to 808 mV at pH 12. This is consistent with the findings in other research, as with the expected pH-dependent shift in electrode potential as well as the characteristics of Ta₂O₅. The Prussian Blue compartment decreased from 676 mV at pH 2 to 199 at pH 12, which can be expected due to the instability of Prussian Blue in alkaline media. The optimization of the electrolyte composition demonstrated that the NaCl concentration had a significant influence on the current density of the cell, up to about 2 M, where the current density increased about 130% from a NaCl concentration of 2 M. This plateau that occurs after 2.0 is where the ionic conductivity benefits maximize before it becomes a transit competition between the ions. The sequential addition of Dy(NO₃)₃ at the optimal NaCl concentration produced a peak of 1.08 mA/cm² at 0.05 M, with the performance of the cell decreasing at higher concentrations.

The fully optimized electrolyte comprising of 0.3 M H₂O₂, 2.0 M NaCl and 0.05 M Dy(NO₃)₃, with a Ta compartment at a pH of 12 and Prussian Blue at a pH of 2, was tested, achieving a OCV of 724 mV and peak power density of 0.283 mW/cm² at 0.80 mA/cm². This is a 40% improvement in the OCV and 172% improvement in the peak power density compared to the unoptimized baseline. The results establish that in a dual configuration, the cell performance can be optimized and yield substantial gains. Future work, as mentioned in the paper, includes using techniques like SEM, EDS and XPS to evaluate the resistance, electrolyte and surface condition of the electrode material, which, as already noted in literature, can be used to evaluate the state of health and degradation of PEMFCs [20]. Longer, sustained operation employing a continuous electrolyte replenishment should be developed to prevent electrolyte self-consumption as well.