Removal of Gaseous Hydrogen Sulfide by a FeOCl/H2O2 Wet Oxidation System

The removal of gaseous hydrogen sulfide using FeOCl/H2O2 was studied. The effects of the FeOCl dosage, the H2O2 concentration, the reaction temperature, and the gas flow rate on the removal of H2S were investigated. The reaction products were analyzed, and the characterization of FeOCl was carried out by X-ray diffraction, scanning electron microscopy, X-ray photoelectron spectroscopy, and electron paramagnetic resonance spectroscopy. Furthermore, radical quenching experiments were carried out using butylated hydroxytoluene, isopropanol, and benzoquinone. It was found that the H2S removal rate for a H2S gas concentration of 160 ppm reached 85.6% when bubbling through 100 mL of an aqueous solution containing FeOCl (1 g/L) and H2O2 (0.33 mol/L) at 293 K with a flow rate of 135 mL/min. Although the dissolution of chlorine in FeOCl was found to result in reduced catalytic performance, the activity was restored after soaking the catalyst in concentrated hydrochloric acid (37%) and subsequent calcination. The mechanism of H2S removal was also discussed, and it was found that this process was controlled by H2S diffusion. FeOCl was found to activate H2O2 and produce radicals, such as •OH and •O2–, resulting in the formation of a water film rich in radicals on the FeOCl surface. Following the diffusion of H2S into the water film, it underwent oxidation by radicals to produce SO42–. Overall, the catalyst and the product can be effectively separated.


INTRODUCTION
Hydrogen sulfide (H 2 S) is a colorless, flammable, and toxic gas with the smell of rotten eggs, and it is a common impurity in natural gas and syngas. Upon its combustion, H 2 S produces sulfur dioxide, which is released into the atmosphere and leads to air pollution. 1 In addition, in industry, the acidity and corrosiveness of H 2 S can cause serious corrosion to equipment and pipelines, while in daily life, an atmospheric concentration of >500 ppm can lead to a partial loss of human neurological function and can even be fatal. 1−4 Thus, to reduce H 2 S emissions, a large number of H 2 S removal technologies have been developed to date, including adsorption/catalytic desulfurization, 5−17 biodesulfurization, 18 membrane separation, 19,20 and washing desulfurization. 21−31 More specifically, the adsorption desulfurization process is simple and the level of investment required is low; however, the regeneration energy consumption is high, and the adsorbent is easily deactivated. 3,4 In contrast, the biological desulfurization technology can operate at room temperature and pressure, its operating costs are low, and it does not tend to cause secondary pollution; however, in the presence of high concentrations of H 2 S, treatment is slow, and the process efficiency tends to be poor. 18 In addition, the membrane separation technology is easily operated and boasts both a high separation efficiency and a low level of pollution; however, the membranes exhibit poor corrosion resistances and short service lives, in addition to being particularly expensive. 19,20 In terms of the chemical washing desulfurization technology, the operation is also simple, and this process exhibits strong adaptability and a high desulfurization rate, resulting in it being the most widely used and mature technology at present. 3,21,23 Nonetheless, the most commonly used traditional absorbents for this process, including ethanolamine and carbonate, require high amounts of energy for their regeneration. 3 In contrast, the wet (solution) oxidation desulfurization technology does not suffer from issues related to absorbent regeneration, the process and equipment required are simple, and a high sulfur recovery tends to be achieved. As a result, this technology has attracted growing attention in recent years. 3,22,24−30 Based on these advantages, Liu et al. [26][27][28][29][30]32 reported the application of the wet oxidation desulfurization technology in an advanced oxidation system. More specifically, hydrogen peroxide, persulfate, and Oxone were used as oxidants, while Fe 3+ , Cu 2+ , Fe 2+ , and UV irradiation were used to promote the production of radicals. Hydrogen peroxide is a low-cost, widely used chemical that does not require complex safety measures. 3 Moreover, Fe 3+ , Cu 2+ , Fe 2+ , and hydrogen peroxide form a homogeneous Fenton or Fenton-like oxidation system, which can effectively remove gaseous H 2 S. 26,27 However, the recovery of Fe 3+ , Cu 2+ , and Fe 2+ is challenging, and during evaporation crystallization, wherein the final oxidation product (sulfuric acid) is recovered by the addition of ammonia to produce ammonium sulfate, 28,30 the product can become contaminated with these metal elements.
Compared with the homogeneous Fenton system, the heterogeneous Fenton system allows the more facile recovery of the catalyst. Among the various catalysts reported to date for this process, iron oxychloride (FeOCl) has been widely studied in terms of the degradation of organic pollutants present in wastewater due to its high catalytic activity, low raw material costs, and its simple preparation and regeneration. 31,33−39 However, to date, the removal of gaseous H 2 S by the FeOCl/ H 2 O 2 system has yet to be examined.
Thus, we, herein, report a novel desulfurization method wherein gaseous H 2 S is removed by bubbling in a FeOCl/ H 2 O 2 wet oxidation system. The technical feasibility of this new technology and the effects of various process parameters (e.g., the FeOCl dosage, the H 2 O 2 concentration, the reaction temperature, and the gas flow rate) on the removal of H 2 S are studied. In addition, the reaction products and the reaction mechanism of H 2 S removal are studied systematically. It is expected that the results of this study will provide the necessary theoretical basis and data for the development of new H 2 S removal technologies and will promote the practical application of the advanced oxidation technology in the field of gaseous H 2 S removal.

MATERIALS AND METHODS
All analytically pure reagents were provided by Aladdin Industrial Corporation (Shanghai, China), and were used directly as received without any additional treatment.

Catalyst Preparation.
FeOCl was prepared by thermal decomposition as described previously. 38,40 More specifically, ferric trichloride (FeCl 3 ·6H 2 O) was ground in an agate mortar, spread flat on the bottom of a crucible (thickness ≤ 3 mm), and heated to 220°C at a rate of 10°C/min in a muffle furnace. After maintaining this temperature for 2.5 h, the crucible was allowed to cool to room temperature, then washed with deionized water, and finally dried overnight at 60°C in a vacuum drying box to obtain the desired FeOCl as a dark fuchsia powder.
2.2. H 2 S Removal. As outlined in Figure 1, the experimental device employed herein consists of a raw material gas supply system, a reaction system, a gas absorption system, and a waste gas treatment system. The raw gas supply system is composed of 300 ppm H 2 S/N 2 (1), high-purity N 2 (2), and a gas mixing system (3−5). The reaction system is mainly composed of three flasks (7) and a heat-collecting magnetic agitator (10), while the gas absorption system to be tested and the waste gas treatment system were prepared using absorption bottles containing zinc acetate (12,14) and sodium hydroxide (13,15), respectively.
The H 2 S/N 2 gas mixture was prepared by mixing the two individual gases as raw materials. By adjusting the flow rates of H 2 S/N 2 and N 2 , the flow rate of the mixture was changed. A heat-collecting magnetic agitator was used to stir the system and to control the temperature. After mixing the gas evenly and adjusting the reaction temperature to the required value, a catalyst adsorption experiment was carried out using only the catalyst and water (100 mL) in a three-neck flask. Following saturation of the catalyst, the desired amount of H 2 O 2 was added to the three-port flask to allow the catalytic oxidation process to take place. The H 2 S concentrations of the inlet and outlet gases were determined by iodometry 41 after being subjected to absorption by zinc acetate. The tail gas was discharged after passing through the waste gas treatment system. The removal rate of H 2 S was calculated by eq 1. During the reaction time of 120 min, four samples were taken at each 30 min interval and the average value was calculated.  In eq 1, η is the removal rate of H 2 S, C in is the H 2 S inlet concentration, and C out is the H 2 S outlet concentration.
2.3. Characterization. The crystal structure of FeOCl was determined by X-ray diffractometry (XRD, DX-2700, China) with Cu Kα radiation. The scanning angle (2θ) was 5−80°and the step size was 0.02°. The morphology of the FeOCl sample was analyzed by scanning electron microscopy (SEM, FEI Inspect F50). X-ray photoelectron spectroscopy (XPS, Thermo Escalab 250Xi) was employed to determine the element valences and chemical information of the FeOCl surface. The binding energy data were calibrated with the C 1s signal at 284.6 eV. Paramagnetic signals were collected on an electron paramagnetic resonance spectrometer (EPR, Brooke A300, Germany) using 5,5-dimethyl-1-pyrrolidine N-oxide (DMPO) to give the EPR spectra of the samples. The oxidation products of H 2 S were analyzed by ion chromatography 32 (IC, ICS-5000).

RESULTS AND DISCUSSION
3.1. Performance of H 2 S Removal by the FeOCl/H 2 O 2 System. The effects of the amount of an FeOCl catalyst, the H 2 O 2 concentration, the reaction temperature, the gas flow rate, and the reaction time on the removal of H 2 S by the FeOCl/H 2 O 2 system were investigated, and the results are presented in Figure 2.
As can be seen from Figure 2A, when the amount of FeOCl is increased from 0 to 1.0 g/L, η increased from 11.1 to 56.6%, indicating that FeOCl has a catalytic effect on the removal of H 2 S by H 2 O 2 , and so η increases with an increase in the FeOCl dosage. This can be accounted for by considering that increasing the amount of FeOCl will provide additional catalytic active centers, 42 which will produce a greater number of active species, and η increases. However, when the amount of FeOCl was increased further to 2 g/L, η increased only from 56.6 to 58.4%. This was likely due to the fact that an excess of the catalyst led to an agglomeration of catalyst particles, thereby hindering the catalytic action of the reaction active sites. 43 As outlined in Figure 2B, when the H 2 O 2 concentration was increased from 0 to 0.33 mol/L, η increased from 0 to 56.6%, although a further increase in the H 2 O 2 concentration to 1.65 mol/L gave only a slight increase in η, i.e., to 60.7%. This was considered that when the concentration of H 2 O 2 in the reaction system is low, the solid catalyst can fully interact with the available H 2 O 2 , and so the amount of an active material (and thus η) increases with an increase in the H 2 O 2 content. 44 However, on increasing hydrogen peroxide continuously, the amount of hydrogen peroxide was in excess to the catalyst relatively, and the reactive material did not increase with the increase in H 2 O 2 ; in addition, adding excess H 2 O 2 would result in the self-consumption of reactive substances. 43 Thus, for the purpose of subsequent experiments, an initial H 2 O 2 concentration of 0.33 mol/L was selected.
As shown in Figure 2C, upon increasing the temperature from 283 to 293 K, η increased from 43.6 to 56.6%, although a further increase in temperature to 313 K resulted in a slight decrease in η (i.e., from 56.6 to 55.4%). In this context, it should be noted that as the temperature increases, the reaction rate also increases initially. 32 However, the solubility of H 2 S in water is reduced at higher temperatures, and so the reduced concentration of reactants ultimately results in a lower reaction rate.
As can be seen from Figure 2D, when the gas flow rate was increased from 135 to 370 mL/min, η decreased from 85.6 to 20.5%. It should be noted that the initial pH of the reaction system is 3.16, and so H 2 S exists mainly in its molecular form in the aqueous environment (K a1 (H 2 S) = 1.3 × 10 −7 , [HS − ]/ [H 2 S] = 1 × 10 −3.56 ). For their effective removal, the H 2 S molecules must diffuse to the catalyst surface, and so as the gas flow rate is increased, the residence time of H 2 S in the reactor decreases, greater volumes of H 2 S are discharged from the reactor prior to diffusion to the catalyst surface, and η is reduced. These results clearly indicate that the removal of H 2 S is controlled by the diffusion of H 2 S.
As outlined in Figure 2E, when the H 2 S gas was continuously introduced into the reaction system over 540 min, η decreased gradually from 56.6 to 40.1%. This can be attributed to a reduced concentration of H 2 O 2 over time and/ or a loss in the catalytic performance of FeOCl. Generally, the oxidation products of H 2 S could be sulfur 32 and/or sulfur oxides (e.g., SO 3 2− and SO 4 2−4 ); in addition, the amount of ineffective H 2 O 2 decomposition is small under acidic conditions (i.e., pH ∼ 3). 45 Assuming that the consumption of H 2 O 2 is mainly attributed to the oxidation of H 2 S, the oxidation of H 2 S to S 0 consumes the least amount of H 2 O 2 (eq 2), 32 while the generation of SO 4 2− consumes the most amount of H 2 O 2 (eq 3). 32 3.2. Analysis of the Reduced Catalytic Performance of FeOCl. In general, the decreased performances of heterogeneous catalysts can be attributed to the reaction products covering or blocking the catalyst active centers or pores or changes in the composition and the structure of the catalyst. 46 As noted above, the products of H 2 S oxidation by H 2 O 2 are sulfur 32 and sulfur oxides. 4 If sulfur is produced, it can cover the active center of the catalyst or block the catalyst pores of the catalyst, whereas if SO 3 2− and SO 4 2− are produced, the pH of the reaction system may decrease. In addition, the interaction between FeOCl and H 2 S may produce FeS. Thus, after the completion of the reaction (i.e., 540 min), the reaction mixture was separated by centrifugation. In the case where SO 3 2− and SO 4 2− are formed, these components should be present in the supernatant, whereas if S 0 and FeS are formed, they would exist in the solid. The liquid phase (or supernatant) was analyzed by ion chromatography and its pH was determined (Table 1); the solid phase (i.e., "used FeOCl") was dried under vacuum at 333 K and characterized by XRD, SEM, and XPS. As listed in Table 1, after 540 min, the pH of the liquid phase was 2.88, and both SO 4 2− and Cl − were detected. Since only FeOCl, H 2 O 2 , H 2 O, and H 2 S/N 2 were added to the reaction system, Cl − present in the liquid phase must originate from FeOCl. The oxidation of H 2 S to SO 4 2− and the dissolution of chlorine from FeOCl to Cl − resulted in the reduced pH of the reaction system (i.e., from 3.16 to 2.88).
It was then considered whether a reduction in the pH of the reaction system would lead to a η decrease. Thus, dilute sulfuric acid was used to regulate the initial pH of the reaction system, and the effect of the initial pH on the removal of H 2 S was investigated ( Figure 3). As shown, when the initial pH values were 2.00, 2.88, and 3.16, the η values were 61.5, 57.2, and 56.6%, respectively, indicating that between values of 2.88 and 3.16, the pH had little effect on the removal of H 2 S. However, when the pH value was 2, η was slightly higher, and this was attributed to the strong acidity of the solution, leading to the dissolution of Fe from the FeOCl, and the amount of H 2 S removed by the homogeneous Fenton system cannot be ignored. 47 These results indicate that a reduction in the system pH during the reaction process does not result in a decrease in η. Figure 4 shows the XRD patterns and SEM images of FeOCl and used FeOCl. As can be seen from Figure 4A, the intensity of the XRD diffraction peaks of used FeOCl decreased significantly compared to FeOCl. In addition, the twodimensional flake morphology of FeOCl can be clearly seen in Figure 4B, 35 while in the case of used FeOCl, this structure was destroyed to a certain extent. Overall, these observations confirm that the crystal structure of FeOCl was partly destroyed during the reaction. Figure 5 shows the XPS spectra of FeOCl, which allowed the elemental composition and valence distribution of the elements on the catalyst surface to be calculated, as listed in Table 2. From Figure 5A, it is apparent that sulfur was not present in used FeOCl, indicating that little or no S 0 and FeS were produced during H 2 S removal. This observation, therefore, indicates that blockage of the catalyst pores or  covering of the catalyst active centers was not responsible for the decreased catalytic activity of FeOCl. As outlined in Table 2, the contents of chlorine in FeOCl and used FeOCl were 26.99 and 8.37% respectively, indicating the loss of chlorine during the reaction. It can be seen from Figure 5B that the O 1s peaks at 530. 35 49 resulting in a decrease in the chlorine content, an increase in the oxygen content, and a higher Fe−O/H 2 O ratio. It is, therefore, possible that the decreased catalytic activity may be caused by the dissolution of chlorine from FeOCl. It was, therefore, expected that the addition of chlorine to used FeOCl could restore its catalytic activity. 50 Thus, used FeOCl was supplemented with chlorine by soaking in a 37% HCl solution (used FeOCl/HCl = 1 g:10 mL) for 30 min, then calcined at 220°C for 1 h, and washed and dried to obtain the regenerated catalyst, i.e., regenerated FeOCl. 50 As shown in Figure 6A, the XRD peak intensity of the regenerated FeOCl is similar to that of FeOCl. Figure 6B shows that η reached 55.4% (average after 120 min) for regenerated FeOCl, which is similar to FeOCl. These results    43 and it is known that H 2 O 2 can approach this polarization plane more easily than water. 31 Thus, the H 2 O 2 molecules adsorbed on the catalyst active sites can be activated to form hydroxyl radicals ( • OH) and superoxide radicals ( • O 2 − ). 32,40,51−54 To confirm the presence of such radicals and their possible contribution to the reaction mechanism, EPR spectroscopy ( Figure 7) and radical quenching experiments (Figure 8) were carried out on the FeOCl/H 2 O 2 system. In the EPR experiment, the DMPO trap reacts with • OH and • O 2 − in solution to form adducts, resulting in the production of characteristic signals. 55 As can be seen from Figure 7 As can be seen from Figure 5C, two large peaks are visible for the Fe 2p orbitals at 711.05 and 724.46 eV, which correspond to the 2p 3/2 and 2p 1/2 binding energies of Fe III , respectively, 56 while two smaller peaks at 712.9 and 726.12 eV correspond to the 2p 3/2 and 2p 1/2 binding energies of Fe II , respectively, 42 To confirm the participation of radicals in the H 2 S removal process, radical quenching experiments were carried out using butylated hydroxytoluene (BHT). 57 As shown in Figure 8C, upon the addition of BHT, η was reduced from 56.6 to 7.8%, which is a similar value to that obtained in the absence of FeOCl (i.e., 10%). These results indicate that the H 2 S removal process did indeed involve participation from radicals. Subsequently, quenching experiments were carried out using isopropanol (IPA) and benzoquinone (BQ). 54 As shown in Figure 8D,E, upon the addition of IPA and BQ, η decreased from 56.6 to 35.4 and 38.5%, respectively, which further indicated that • OH and • O 2 − were involved in the process of H 2 S removal by the FeOCl/H 2 O 2 system. The chemical reactions involved in this process are given by eqs 8−10, 32 Upon the simultaneous addition of both IPA and BQ ( Figure 8F), η was 24.2%, which is higher than that obtained when BHT was added (i.e., 7.8% η), indicating that other radicals were also involved in the H 2 S removal process. As can be seen from eq 5, HO 2 • is also produced during the catalytic process, but this species is further decomposed to produce • O 2 − (eq 6), and the oxidizability of HO 2 • (oxidation potential 1.51) is significantly weaker than that of • O 2 − (oxidation potential 2.21) and • OH (oxidation potential 2.80). 59 In addition to • OH, • O 2 − , and HO 2

•
, there are other forms of radicals involved in the reaction.
Wang 60 found that the radicals produced by H 2 O 2 activation can be divided into two states, namely adsorbed and free states. IPA is a hydrophilic radical scavenger, which cannot effectively quench adsorbed radicals; 61 BQ is also hydrophilic and it was therefore considered that BQ can only quench free superoxide radicals. Radical quenching experiments show that free radicals are involved in the removal of H 2 S. According to the process of the catalytic reaction, eqs 4 and 13−15 can be given as follows In eqs 13−15, the subscript "ads" denotes radicals in the adsorbed state and "free" denotes radicals in the free state.
According to eqs 4 and 13−15, there are adsorbed radicals in the system. Qian 55 reported that the adsorbed radicals are easy to oxidize the pollutants adsorbed on the catalyst. If H 2 S in the solution can be adsorbed on FeOCl, it shows that the adsorbed radicals are likely to participate in the oxidation of H 2 S. Thus, the adsorption of H 2 S by FeOCl in the aqueous phase was investigated and the adsorption kinetics was analyzed, as shown in Figure 9 and Table 3.
It can be seen from Figure 9A that with an increase in temperature, the solubility of H 2 S in solution decreases, but the adsorption capacity of H 2 S increases. Obviously, the adsorption of H 2 S by FeOCl in solution accords with the characteristics of chemical adsorption. 62 As indicated in the fitting results shown in Figure 9B,C and Table 3, the R 2 of the  pseudo-second-order adsorption kinetics of H 2 S on FeOCl is larger than that of the pseudo-first-order adsorption kinetics at 283, 293, and 303 K, indicating that the pseudo-second-order model can better describe the adsorption process. 2,62 According to the results of kinetic analysis, enthalpy (ΔH) of H 2 S adsorption by FeOCl can be calculated by eqs 16−18 to be 21.5 kJ/mol. Generally, ΔH of the chemisorption is >40 kJ/mol, 63 which indicates that the interaction between FeOCl and H 2 S is relatively weak, and H 2 S could be activated by adsorbing. Therefore, we speculate that the adsorbed radicals were involved in the removal of H 2 S. This may also be the reason that FeS was not generated during the catalysis and sulfur was not detected in the XPS.
In eqs 16−18, K is the adsorption equilibrium constant and q e is the equilibrium adsorption capacity of H 2 S (saturated adsorption capacity, mg/g). According to the results of kinetic analysis, q e = q e2 and c e is the concentration of H 2 S in the aqueous solution at the point of adsorption equilibrium (mg/ g). Since the concentration of H 2 S in the gas phase remains constant (160 ppm) during the adsorption experiment, the value of c e can be obtained according to Henry's law. ΔG is the free energy (kJ/mol), ΔH is the enthalpy (kJ/mol), ΔS is the entropy (J/(mol·K)), R is the gas constant (8.314 J/(mol·K)), and T is the adsorption temperature (K).
3.4. Mechanistic Analysis. In general, the lifetimes of radicals are extremely short, and so they are mainly concentrated in the water film close to the surface of the solid catalyst. 34 In heterogeneous catalytic reactions, the reaction process is generally composed of several processes occurring in series, such as diffusion of the H 2 S to the surface of the catalyst, the adsorption of H 2 S onto the catalyst surface, and conversion of the adsorbed H 2 S. Due to the fact that a water film rich in radicals is present on the surface of FeOCl, and the free radicals can react directly with H 2 S, 32 it is possible that two reaction paths exist for H 2 S: (I) the adsorbed H 2 S is oxidized by adsorbed radicals and (II) the not adsorbed H 2 S is oxidized by free radicals. Based on the abovementioned analysis, a mechanism was established for H 2 S removal by the FeOCl/H 2 O 2 system, as shown in Figure 10. The oxidation product is mainly SO 4 2− , and the diffusion of H 2 S controls the rate of H 2 S removal.
The resulting SO 4 2− can be recovered by the addition of ammonia to produce ammonium sulfate, 32 thereby allowing the effective recovery of the reaction products to achieve zero discharge of liquid waste from our proposed system.

CONCLUSIONS
This paper reports the study on the performance of gaseous H 2 S removal by the FeOCl/H 2 O 2 system. It was found that upon increasing the amount of a catalyst, the H 2 O 2 concentration, and the reaction temperature, the removal efficiency of H 2 S initially increased prior to becoming stable. In contrast, upon increasing the gas flow rate, the removal efficiency of H 2 S decreased, and so it was apparent that the process of H 2 S removal by FeOCl/H 2 O 2 is a diffusioncontrolled process. The removal rate of 160 ppm H 2 S gas reached 85.6% when bubbling through 100 mL of an aqueous solution containing FeOCl (1 g/L) and H 2 O 2 (0.33 mol/L) at 293 K with a flow rate of 135 mL/min. Although the dissolution of chlorine in FeOCl caused a reduction in the catalytic performance, the catalytic activity could be restored by soaking in hydrochloric acid. The main active species involved in the removal of H 2 S were found to be radicals, such as • OH and • O 2 − , which were produced from H 2 O 2 upon catalysis by FeOCl. These radicals were likely present in the adsorbed and free states, and the free radicals were mainly distributed in the water film close to the FeOCl surface. In terms of the process mechanism, H 2 S initially diffuses toward the vicinity of FeOCl or adsorbs onto the catalyst surface. Subsequently, H 2 S is oxidized by free and adsorbed radicals, Figure 10. Schematic representation of the heterogeneous degradation of H 2 S by the FeOCl/H 2 O 2 system. Path I, the adsorbed H 2 S is oxidized by adsorbed radicals, and path II, the unadsorbed H 2 S is oxidized by free radicals. respectively, to give SO 4 2− . Importantly, the catalyst and the product of this system can be separated effectively. To enhance the removal rate of H 2 S further, it is necessary to increase the residence time of H 2 S in the reactor, such as replacing the three-neck flask with a tower reactor. Overall, this process was found to be a simple and efficient method for H 2 S purification, and following scale-up, could be suitable for industrial applications.