## 1 Introduction  
The original simple fuel cell was invented by [PERSON] in 1894, who thought that instead of typical combustion, certain fuels could undergo oxidation through controlled electromechanical mechanisms [1]. It was not until the mid-1980s that fuel cells were given formal and practical consideration for generating on-demand electrical power in remote areas. Today, fuel cells are a renowned category with many industrial applications and substantial research in various types and operations. A typical fuel cell contains two electrodes, a cathode and an anode, in two segments isolated by a proton conductive barrier material. The fuel flows through the anode compartment while the reactant flows through the cathode compartment. A redox reaction occurs facilitated by the passage of protons across the proton exchange barrier [2]. The chemical reaction creates a net electrical potential difference at the cathode and anode sites. Hydrogen is a popular type of fuel cell where hydrogen gas is reacted with oxygen; however, the use of hydrogen peroxide as fuel has promising capabilities to improve upon hydrogen fuel cells. Hydrogen peroxide is stable in liquid form, making it easier to store and transport without the need for high-pressure methods like with hydrogen. Hydrogen peroxide, even at very high concentrations, can be reliably stored in containers made of passivated resistant stainless-steel, high-density polyethylene (HDPE), or glass. Pressurized hydrogen will almost always require the use of a reinforced metallic capsule, and despite all efforts, hydrogen has inherent material damaging risks. A concern with the use of hydrogen peroxide as both the fuel and oxidant pertains to how it can undergo a mixed potential reaction calledhydrogen peroxide disproportionation. In this mixed reaction, oxidation and reduction can occur between the peroxide molecules without reacting with the electrode. This causes the degradation of the peroxide without the formation of any potential difference across the electrodes. The decomposition of Hydrogen peroxide in acidic solutions occurs through the redox reactions shown in Equations (1) and (2) [3]. In an alkaline solution, the redox reaction differs as the peroxide enters the reaction in its dissociated form, as shown in Equations (3) and (4). In these reactions, the hydrogen peroxide reacts with the anode to produce hydrogen protons and electrons. The electrons travel from the anode to the cathode through a circuit to generate electricity. The electrons then react with the peroxide and hydrogen protons at the cathode electrode to produce water. The net cell electrochemical reaction of the hydrogen peroxide, as shown in Equation (5), generates a predicted electrical potential of 1.09 V. This theoretical voltage is comparable to other common fuel cells such as hydrogen fuel cells (1.23 V) and methanol fuel cells (1.21 V). \[\rm{H_{2}O_{2}}\longrightarrow\rm{O_{2}}+2\,H^{+}+2\,e^{-} \tag{1}\]  
\[\rm{H_{2}O_{2}}+2\,H^{+}+2\,e^{-}\longrightarrow 2\,H_{2}O \tag{2}\]  
\[\rm{HO_{2}}^{-}+\rm{HO^{-}}\longrightarrow\rm{O_{2}}+H_{2}O+2\,e^{-} \tag{3}\]  
\[\rm{HO_{2}}^{-}+H_{2}O+2\,e^{-}\longrightarrow 3\,OH^{-} \tag{4}\]  
\[2\,H_{2}O_{2}+O_{2}\longrightarrow 2\,H_{2}O \tag{5}\]  
[PERSON] et al. conducted the first academic investigation of a single-compartment hydrogen peroxide, published in 2008 [4]. The study assessed the effectiveness of Gold (Au), Platinum (Pt), Palladium (Pd), Silver (Ag), and Nickel (Ni) as electrode materials using alkaline hydrogen peroxide as the common electrolyte. Since then, there have been various studies investigating hydrogen peroxide fuel cells as shown in Table 1.