One of the crucial steps in direct electrolysis of H₂S at low temperatures is their absorption by a suitable electrolyte. This fluid should be chemically stable and thermodynamically inactive within the electrolysis potential difference. Early works demonstrated the possibility of using alkaline solutions for this purpose, particularly solutions containing NaOH (Anani et al., 1990; Mao et al., 1991). In those works, attention was shifted from the material composition of the SOR electrocatalyst, as their primary objective was to find the optimal H₂S processing conditions. This involves splitting H₂S with minimal sulfur deposition on the electrocatalysts at ∼80°C. Anani et al. (Anani et al., 1990) used graphite, nickel, nickel-chromium, and titanium as SOR electrodes. However, they did not present a meaningful comparison of their electrocatalytic activity as they only mentioned a worse passivation effect in the metallic electrodes. Indeed, at that time, the deactivation of electrocatalysts by sulfur precipitation during the electrolysis of polysulfide solutions was a well-known problem (Fetzer, 1928; Hodes et al., 1980; Buckley et al., 1987), and several attempts to palliate it are documented, including the use of organic solvents to remove it from the electrocatalyst (Shih and Lee, 1986).

Research towards optimizing the H₂S splitting continued, looking for better SOR electrocatalysts and more practical methods to avoid their deactivation. Subsequent works expanded the list of SOR electrocatalytic materials, adding LaSrMnO₃ and Raney-nickel. Petrov and Srinivasan (Petrov and Srinivasan, 1996) demonstrated that LaSrMnO₃ had better SOR electrocatalytic performance than graphite, CoS, Raney-nickel, and Pt/C electrodes. They reached a current density of 300 mAcm⁻² at a cell potential of only 1 V. Their process scrubbed H₂S in a NaOH solution, then alkalinized the solution to pH 14, and electrolyzed it. All those steps were done in separate chambers, as Anani et al. (Anani et al., 1990) had previously proposed. Petrov and Srinivasan improved Anani et al.‘s method by doing the final sulfur precipitation in a separate chamber, preventing sulfur precipitation in the graphite, CoS, and LaSrMnO₃ electrodes but not in the Raney-nickel, Ni, and Pt/C electrodes.

Concomitantly, other researchers focused in understanding the reaction mechanism of sulfide electrooxidation. Cobalt phthalocyanine traces, used previously to catalyze the electroreduction of oxygen in KNO₃ aqueous solutions (Komorsky-Lovrić, 1995), were studied as electrocatalysts for the oxidation of sulfide ions by Komorsky-Lovrić et al. (Komorsky-Lovrić et al., 1997) via cyclic voltammetry (CV). Under the chemical conditions of their experiments, the CV measurements were consistent with a two-step oxidation of HS⁻ ions: 2 H S − + O H − + K + → K H S 2 a d s + 2 e − + H 2 O (2) 2 H S − + O H − → H S 2 − + 2 e − + H 2 O (3)

In the case of a low sulfide ion concentration (<10⁻⁴ M), the KHS₂ adsorbed species would dissociate following the reaction shown in Eq. 4. K H S 2 a d s → S a d s + H S − + K + (4)

Komorsky-Lovrić et al. (1997) stated that the adsorbed sulfur stays on the surface of the electrocatalyst and works as a bridge for the charge transfer process between the electrode and HS⁻ ions. However, it is essential to note that the proposed reaction mechanism does not yield free elemental sulfur as a final byproduct and that there were no analytical measurements of the species in the electrolyte after the redox reaction. Moreover, there was no description of gas evolution in electrodes, so it is unclear whether cobalt phthalocyanine-based SOR electrodes can perform the corresponding electrooxidation reaction to produce elemental sulfur.

Chen and Miller (Chen and Miller, 2004; Miller and Chen, 2005) performed further studies related to the mechanism of the electrochemical oxidation of sulfide ions. The SOR electrode used in their experiments was Ti/Ta₂O₅-IrO₂, and they performed different electrochemical characterizations under different conditions, including variations in temperature and sulfide concentration in the aqueous electrolyte. In their experiments, temporal oscillations of the SOR electrode potential (vs. a reference electrode) were present in specific galvanostatic conditions and at specific current intervals during linear galvanic voltammograms. The analysis of their chronopotentiometry observations indicates that the origin of the oscillations was the alternation of the dominant oxidation reaction performed by the electrode. In the lower part of the potential oscillation, the dominant reaction is the oxidation of sulfide into sulfur, causing the sulfur deposition on the electrode and therefore inducing an increase of the electrode potential to keep the current at the set value. However, this increment in potential favors the oxygen evolution reaction (OER) from the aqueous electrolyte, making water oxidation the dominant reaction in the oscillation peak potential. The shift towards OER increases the availability of HS⁻ and S^2- ions as sulfur oxidation is minimized, increasing the dissolution of sulfur through the reactions displayed in Eq. 5 and Eq. 6. S + H S − + O H − → S 2 2 − + H 2 O (5) S + S 2 − → S 3 2 − (6)

The enhanced sulfur solubility and the increased mass transfer rate generated by the oxygen evolution led to the removal of sulfur deposits from the SOR electrode, promoting the rebound of the electrode potential at the lower bound of the oscillation. The continuous formation and removal of sulfur deposits on the surface of the electrode were observed during those measurements. This further supports the proposed reaction mechanism. These results highlight the importance of the selection of electrode materials, electrolytes, and electrical conditions of the electrolysis. The latter encompasses current density and potential, which influence the preferred chemical reactions. Such observation becomes even more relevant when there is great interest in applying these electrolytic processes at an industrial scale, where it is necessary to use high currents and voltages to obtain a higher yield of desired products.

Subsequent works explored the possibility of using electrolysis for H₂S splitting in applications beyond the discussed cases of fossil fuel and natural gas scenarios. Sanli et al. (Sanli et al., 2014) researched the attainability of oxidizing sulfide ions contained in artificial seawater with chemical conditions like the ones found in the Black Sea. Their investigations used MoS₂ as the SOR electrode, which exhibited electrocatalytic activity towards the oxidation of HS⁻ into elemental sulfur. However, the final purpose of their work was to use the SOR electrode in H₂S fuel cells, where the catalytic HS⁻ oxidation must occur spontaneously to generate electricity besides elemental sulfur. The idea of oxidizing sulfide from brines was also explored by Ateya et al. (Ateya et al., 2005). They used carbon-felt SOR electrodes to drive H₂S splitting from wastewater brines. Alternatively, Karapekmez and Dincer (Karapekmez and Dincer, 2018) presented an H₂S abatement system using a proton exchange membrane electrolyzer (PEME). The H₂S PEME had a CsHSO₄ electrolyte and operated at 150°C. Their system was designed to treat the H₂S generated in geothermal power plants, and its practical implementation led to a substantial decrease in H₂S emissions.

The passivation of SOR electrodes had been the bottleneck in the development of low-temperature H₂S splitting systems. Recently, two research approaches, shown schematically in Figure 1, demonstrated the possibility of avoiding the passivation of SOR electrodes, one through the design of nanostructured electrocatalysts and the other using a novel organic sulfur solvent. In the case of nanostructured electrodes, Zhang et al. (Zhang et al., 2020) reported the fabrication of a SOR electrode consisting of CoNi nanoparticles encapsulated in graphene shells and supported in Ni foam. Three-electrode measurements showed that using this electrode, the onset potential of sulfide oxidation was only ∼0.25 V vs. reversible hydrogen electrode (RHE), 1.24 V lower than the onset potential of water oxidation. Additionally, the electrode exhibited a stable performance during a 500 h chronoamperometry test using a Na₂S/NaOH aqueous solution with a current density of around 30 mAcm⁻². These results showcase the possibility of generating H₂ reliably, with a Faradaic efficiency of about 98% and using significantly less energy than water splitting. The chronopotentiometry analysis of the nanostructured electrode in more realistic conditions was also reported, where it was tested for 1,200 h in a 1 M NaOH solution saturated with 2% H₂S/Syngas (50% CO and 50% H₂) at a current density of 20 mAcm⁻². In such conditions, the electrode’s potential was maintained at around 0.5 V vs. RHE with minimal variations. Density functional theory calculations showed that the free adsorption energy of the sulfide oxidation intermediate is minimized in graphene-encapsulated CoNi nanoparticles compared with pure graphene or CoNi nanoparticles alone. This characteristic is generally accepted as an indicator of good catalytic material (Nørskov et al., 2009), thus explaining the excellent activity of the electrode towards the SOR and its capacity to avoid sulfur deposition on its surface.

Nanostructured electrodes exhibited further advantages in the work led by Kumar and Nagaiah (Kumar and Nagaiah, 2022). They used CoFeS₂ nanograins conjugated with nitrogen-containing carbon as a SOR electrode. They tested several stoichiometries of the sulfide material, finding a better electrocatalytic activity towards the SOR using an electrode with a Co:Fe ratio of 3:1. Linear sweep voltammograms of this electrode showed a SOR onset potential of just 0.23 V vs. RHE. The current of the electrode in chronopotentiometry measurements for 120 h was about 55 mAcm⁻² at only 0.3 V vs. RHE, a better performance than the nanostructured SOR electrode from the work of Zhang et al. (Zhang et al., 2020) with a comparable Faradaic efficiency for H₂ generation (∼97%). The electrode’s current during the 120-h characterization was relatively stable, with a slight decrease over time due to a reduction of the electrochemically active area to 81% of its initial value, demonstrating a minor passivation effect suffered by the electrode.

Ma et al. (2020) proposed another approach to suppress the passivation of the SOR electrode. They electrolyzed H₂S in a mixture of an ionic liquid (IL) [C₃OHmim]- BF₄, tetraethylene glycol dimethyl ether (TGDE), and monoethanolamide (MEA). The IL, TGDE, and MEA mixture exhibited a superior sulfur solubility above 20°C when compared with a mixture of IL and TGDE. This is ascribed to the protonation of MEA by H₂S, as indicated in Eq. 7. M E A + H 2 S → M E A H + + H S − (7)

They took advantage of the gradient of the sulfur solubility as a function of temperature to propose a cyclic electrolytic process. A Pt microdisk was used as a SOR electrode and the temperature was modified at different stages of the cycle. During H₂S electrolysis, the temperature of the electrolyte is kept at 50°C. As the electrolysis process continues, the electrolyte’s color shifts from transparent to yellow, indicating the presence of elemental sulfur dissolved in the electrolyte. The electrolysis is stopped, and the electrode is removed before the electrolyte saturates with dissolved sulfur to avoid its precipitation on the surface of the Pt microdisk. Next, the electrolyte is cooled down to room temperature, which decreases its sulfur solubility, enhancing the precipitation of sulfur. Then the solid sulfur is separated from the electrolyte, which can be used to absorb H₂S again and start a new electrolytic cycle. The electrode’s endurance test consisted of three 7-h electrolytic cycles with a current density in the range of 10–13 mAcm⁻². The Faraday efficiency of the H₂ evolution presented a slight decrease from 90.5% to 89.3% from the first to the third electrolytic cycles. These relatively low values were attributed to the formation of polysulfides. The “shuttling mechanism,” could explain the large overpotential (∼1.63 V at ∼10.5 mAcm⁻²) in comparison to the nanostructured electrode proposed by Zhang et al. (0.4 V at ∼30 mAcm⁻²) (Zhang et al., 2020). The report does not discuss the origin of the large overpotential or if it originates in the oxidation or reduction side. However, the cyclic voltammograms suggest that the electrocatalytic activity of the Pt microdisk is larger when using the electrolyte containing MEA. The origin of the overpotential is likely related to the slower kinetics of this amine, for example, in the reduction of the protonated MEA species formed by the reaction shown in Eq. 7.