## 2 Experimental Part  
### Methodology  
In the case under consideration, i.e., where in situ energy production from the hydrogen sulfate present in marine water is anticipated, the following restrictions were kept in mind:  
* No discharge of any toxic or environmentally unfriendly waste into the sea. Therefore, no use of chemical oxidizers for the cathode processes (e.g., hexacyanoferrate, etc.) instead of oxygen is admissible;
* No use of any other supporting electrolytes but marine water (no sodium or potassium hydroxides);
* Use of HS- or S2- as substrates only;
* No metal parts in the fuel cell design;
* Metal electrodes are undesired. Figure 1: Working principle of a sulfide-driven fuel cell associated with proton (**a**) and OH- (**b**) exchange across the membrane. The experiments consisted in parallel measurements of the cell potential (in an open circuit and with a load) and the current in specially designed sulfide-driven fuel cells under batch and continuous conditions. From these results voltammetric polarization curves were constructed. From the slope of the linear part of the curves the cell Ohmic resistance was determined and subsequently the power density was calculated. Afterwards, cell discharge through a selected Ohmic resistance in batch or continuous mode was accomplished, accompanied by sampling from the inlet and outlet streams in the continuous experiments or periodically during the batch ones. Electric current and cell potentials were measured and monitored continuously during all of the runs. The results of the analyses for the sulfide electrochemical oxidation were compared to the ones calculated from the current values according to [PERSON]'s law:  
\[\frac{m}{t}=\frac{Mi}{nF},m=\frac{M}{nF}\int_{0}^{t}idt \tag{1}\]  
where:  
* \(i\)--electric current, A;
* \(m\)--mass of reacting substance, g;
* \(t\)--time, s;
* \(M\)--molar mass of reacting substance, g;
* \(n\)--number of exchanged electrons;
* \(F\) = 96,484 C mol\({}^{-1}\), Faraday constant. One must have in mind that there are many sulfide oxidation reactions incorporating one to eight exchanged electrons. Short excerpt of a list of sulfide redox reactions [27] relevant to the studied process is shown in Table 3. Both proton and hydroxide ion exchange processes are possible. Other reactions are also possible with polysulfide ions as an end-product. The latter are observed at higher sulfide concentrations over a certain pH range, cf. Figure 2. One can see that over the pH range of 7-9 hydrosulfide anions are predominant. Such is the case in the Black Sea water. \begin{table}
\begin{tabular}{c c c} \hline \hline
**Reversible Redox Anode Reactions** & **Number of Exchanged** & **Standard Electrode** \\ **(Short excerpt)** & **Electrons** & **Potential [V], 25 \({}^{\circ}\)C** \\ \hline \({\rm SO_{4}}^{2-}+{\rm H_{2}O}+2e={\rm SO_{3}}^{2-}+2{\rm OH}^{-}\) & 2 & \(-0.93\) \\ \({\rm SO_{3}}^{2-}+3{\rm H_{2}O}+6e={\rm S}^{2-}+6{\rm OH}^{-}\) & 6 & \(-0.66\) \\ \({\rm S_{2}}^{2-}+2e=2{\rm S}^{2-}\) & 1 & \(-0.524\) \\ \({\rm S+2}e={\rm S}^{2-}\) & 2 & \(-0.48\) \\ \(2{\rm SO_{4}}^{2-}+4{\rm H}^{+}+2e={\rm S_{2}}{\rm O_{6}}^{2-}+2{\rm H_{2}O}\) & 2 & \(-0.22\) \\ \({\rm S+H}^{+}+2e={\rm HS}^{-}\) & 2 & \(-0.065\) \\ \({\rm S_{2}O_{3}}^{2-}+6{\rm H}^{+}+8e=2{\rm S}^{2-}+3{\rm H_{2}O}\) & 4 & \(-0.006\) \\ \({\rm HSO_{3}}^{-}+5{\rm H}^{+}+4e=5+3{\rm H_{2}O}\) & 4 & \(0\) \\ \({\rm SO_{4}}^{2-}+8{\rm H}^{+}+8e=5{\rm S}^{2-}+4{\rm H_{2}O}\) & 8 & \(0.149\) \\ \({\rm S_{2}O_{3}}^{2-}+8{\rm H}^{+}+8e=2{\rm HS}^{-}+3{\rm H_{2}O}\) & 4 & \(0,2\) \\ \hline \hline \end{tabular}
\end{table} 
Table 3: Short excerpt of sulfide oxidation reactions, from Reference [27]. Among the various redox reactions, in which sulfur-containing compounds and ions are involved, is the formation and oxidation of thiosulfate and other thio-anions, cf. Table 3. Thiosulfate is stable in neutral and alkaline media, as also demonstrated in our case.