Introduction to the desulfurization process using hydrogen sulfide-complexed iron for dry gas treatment

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The complexed iron-based desulfurization process uses iron ions as a catalyst to directly convert hydrogen sulfide into elemental sulfur through liquid-phase redox reactions, achieving a hydrogen sulfide removal rate of over 99%. The main advantages of this new desulfurization process are its simplicity of operation, high sulfur capacity, and environmental friendliness. It can be used to remove hydrogen sulfide gas from any type of gas stream, such as refinery gas, natural gas, sulfur-containing gases, Claus off-gases, amine-process gases, and Selexol off-gases. The desulfurization efficiency is not affected by the concentrations of carbon dioxide, sulfur, or hydrogen sulfide in the gas stream. It is the preferred process for treating sulfur-containing gases with a sulfur output of 0.2–20 tons per day. The process principle is as follows: This process improves the Claus reaction, providing a constant-temperature, low-cost method for removing hydrogen sulfide: H2S + 1/2 O2 → H2O + S (1). This reaction takes place in an aqueous solution, and it is facilitated by the addition of water-soluble metal ions. Because in air or acidic gases, water-soluble metal ions are easily oxidized by oxygen, and a stable electrode is used to oxidize sulfide ions into elemental sulfur. In other words, since the reaction takes place in an aqueous solution, this aqueous solution containing metal ions can easily release the electrons of divalent sulfur (HS-) to form elemental sulfur, and transfer those electrons to oxygen during the regeneration process. Although many metal ions can carry out the above reaction, iron ions are chosen for this process due to their low cost and non-toxicity. The basic chemical reactions in the desulfurization process can be divided into two parts: absorption and regeneration, as follows:
Absorption part: The prepared solution absorbs H2S gas:
H2S (gaseous hydrogen sulfide) + H2O (liquid water) ↔ H2S (aqueous hydrogen sulfide) + H2O (water) (2)
Ionization: H2S (aqueous hydrogen sulfide) ↔ H+ + HS- (3)
Higher-valent iron ions (Fe3+) oxidize divalent sulfur: HS- + 2Fe3+ → 2Fe2+ + H+ + S0 (4)
Overall equation for the absorption part (combination of equations 2, 3, and 4):
H2S (gaseous hydrogen sulfide) + 2Fe3+ → 2H+ + S° + 2Fe2+ (5)
Regeneration part: The prepared solution absorbs oxygen:
1/2 O2 (gas) + H2O ↔ 1/2 O2 (liquid) + H2O (6)
Regeneration reaction of ferrous ions (Fe2+):
1/2 O2 (liquid) + H2O + 2Fe2+ → 2OH- + 2Fe3+ (7)
Overall equation for the regeneration part (combination of equations 6 and 7):
1/2 O2 (gas) + H2O + 2Fe2+ → 2OH- + 2Fe3+ (8)
Combining equations 5 and 8 yields the improved Claus reaction equation, as shown in (1). In the overall reaction, the role of iron ions is to release the electrons generated in the absorption reaction into the regeneration reaction. Since the formation of each molecule of elemental sulfur requires the consumption of two iron atoms, at least two iron atoms must be provided during the reaction process. Thus, iron ions act as reactants. However, iron ions are not consumed in the overall reaction; they act as a catalyst for the reaction between hydrogen sulfide and oxygen. Due to this dual function, iron ion complexes are generally defined as catalytic reactants. In the aqueous solution, neither ferrous ions (Fe2+) nor ferric ions (Fe3+) can exist stably; under normal conditions, they tend to form iron oxide or iron sulfide precipitates through the following reactions: Fe3+ + 3OH- → Fe(OH)3 (solid) (9) Fe2+ + S2- → FeS (solid) (10) To prevent the formation of such precipitates, a chelating agent is used in the system, allowing the iron ions to remain stable in the solution over a wide range of pH values.