Effect of Water Content on Properties of Homogeneous [bmim]Fe(III)Cl₄–H₂O Mixtures and Their Application in Oxidative Absorption of H₂S

Abstract

The potential of 1-butyl-3-methylimidazolium tetrachloroferrate ([bmim]Fe(III)Cl₄) for replacing an iron(III) chelate catalytic solution in the catalytic oxidation of H₂S is attributed to its no side reaction and no degradation of the chelating agent. The catalytic oxidation product of water in non-aqueous [bmim]Fe(III)Cl₄ possibly has an influence on the oxidative absorption of H₂S. Water and hydrophobic [bmim]Fe(III)Cl₄ mixtures at water volume percents from 40% to 70% formed separate phases after srirring, without affecting the oxidative absorption of hydrogen sulfide. Then, studies on the properties of homogeneous [bmim]Fe(III)Cl₄–H₂O mixtures at water volume percents in the range of 5.88–30% and above 80% reveal that these mixtures are both Brønsted and Lewis acids at vol % (H₂O) ≤ 30%, and only Lewis acids at vol % (H₂O) ≥ 80%. Raman spectra showed that [bmim]Fe(III)Cl₄ was the dominating species at vol % (H₂O) ≤ 30%, in contrast, [bmim]Fe(III)Cl₄ decomposed into FeCl₃·2H₂O and [bmim]Cl at vol % (H₂O) ≥ 80%. Further research on oxidative absorption of H₂S by homogeneous [bmim]Fe(III)Cl₄–H₂O mixtures demonstrated that [bmim]Fe(III)Cl₄ was reduced by H₂S to [bmim]Fe(II)Cl₄H and FeCl₃·2H₂O was reduced to FeCl₂, at the same time, H₂S was oxidized to S₈. In addition, the decrease in acidity caused by increasing the water content increased the weight percent of absorbed H₂S, and decreased volatile HCl emissions. However, it is difficult to prevent the suspended S₈ generated at vol % (H₂O) ≥ 80% from the formation of sulfur blockage. Therefore, oxidative absorption of H₂S by [bmim]Fe(III)Cl₄–H₂O mixtures is feasible at vol % (H₂O) < 80% without sulfur blockage.

1. Introduction

[bmim]Fe(III)Cl₄ (1-butyl-3-methylimidazolium tetrachloroferrate) is an ionic liquid with unusual properties, including low volatility, high thermal stability, high hydrophobicity [1], low viscosity [2], and low surface tension [3]. [bmim]Fe(III)Cl₄ has also been found to be a potentially suitable solvent for gas separation and a highly effective catalyst for many reactions [4,5,6].

There has been a considerable recent interest in the use of [bmim]Fe(III)Cl₄ as an alternative to an iron(III) chelating solution for catalytic oxidation of H₂S. Iron(III) chelating solutions have several disadvantages that are associated with liquid phase oxidation processes, including degradation of the chelating agent [7] and sulfur oxo-acid salt formation [8]. Liquid phase oxidation processes using [bmim]Fe(III)Cl₄, however, do not require a complexing agent. In addition, sulfur oxo-acid salts cannot form in acidic media [9]. Catalytic-oxidation of H₂S by [bmim]Fe(III)Cl₄ differs from other heterogeneous catalytic reactions in that it is comprised of two simultaneous reactions that occur in the same or a separate vessel [10].

A description of these reactions follows. [bmim]Fe(III)Cl₄ is reduced by H₂S and then oxidatively regenerated by air or oxygen. The two steps demonstrate that [bmim]Fe(III)Cl₄ is both a solvent and a catalyst. The overall reaction is shown in Equation (1).

H₂S + 0.5O₂ → S + H₂O

(1)

Apart from sulfur, which is easily recovered by filtration, only H₂O is generated. Since [bmim]Fe(III)Cl₄ is hydrophobic [1], mixtures of [bmim]Fe(III)Cl₄ and water should form two phases that should be easily separable. However, Lee et al. reported that 20% (v/v) of [bmim]Fe(III)Cl₄ and water were fully miscible after vigorous shaking [11]. Wang et al. also reported that a [bmim]Fe(III)Cl₄ can be dissolved in 80% (v/v) H₂O, while it was difficult for even a spot of H₂O to be dissolved in pure [bmim]Fe(III)Cl₄ [12]. Additional studies indicated that the presence of water in ionic liquids that have no metal affected their physical and chemical properties. The surface tension of ammonium ionic liquids in aqueous solutions increased non-linearly with water content [13]. The viscosity of hydrophilic 1-butyl-3-methylimidazolium-based ionic liquids also strongly depends on water content [14]. In addition, high hydration numbers are observed for imidazolium-based ionic liquids in aqueous mixtures [15]. Furthermore, solute-induced dissolution of hydrophobic ionic liquids in water was observed by Rickert et al. [16]. Nevertheless, to our knowledge, there are no literature reports on the effect of water on the properties of metal-containing ionic liquids, such as [bmim]Fe(III)Cl₄. Since aqueous mixtures of [bmim]Cl influence the solution structures of two archetype proteins [17], and imidazolium perrhenate ionic liquids in aqueous hydrogen peroxide were proved to be efficient catalysts for the selective oxidation of sulfides to sulfones [18], it is possible that water in [bmim]Fe(III)Cl₄ has an influence on the oxidative absorption of H₂S.

In this paper, we report our findings on the properties of homogeneous mixtures of [bmim]Fe(III)Cl₄ and H₂O, and their effect on the oxidative absorption of H₂S. Additional work on the regeneration of [bmim]Fe(III)Cl₄ from these aqueous homogeneous mixtures will be discussed in a future paper.

2. Results and Discussion

2.1. Effect of Water Concentration on the Formation of [bmim]Fe(III)Cl₄–H₂O Mixtures

The effect of water concentration on the formation of homogeneous [bmim]Fe(III)Cl₄–H₂O mixtures was studied by adding different amounts of water to [bmim]Fe(III)Cl₄.

Figure 1 shows that dark brown homogeneous liquids are obtained at x(H₂O) values from 5.88% to 30%. At x(H₂O) values from 40% to 70%, however, the mixtures form two phases. The upper aqueous phase is brownish yellow and the lower oil phase is dark brown. The homogeneous liquids at x(H₂O) ≥ 80% are brownish yellow. Since there wasn’t the presence of structured water in the hydrophobic ionic liquid phase [19], it is very easy to separate water from [bmim]Fe(III)Cl₄ in the two-phase mixtures without affecting the oxidative absorption of hydrogen sulphide. However, it is difficult to separate water from the homogeneous [bmim]Fe(III)Cl₄–H₂O mixtures without evaporation method [11,12]. On the one hand, the acidity of homogeneous [bmim]Fe(III)Cl₄–H₂O mixtures may change with the increment of water. On the other hand, as the x(H₂O) in homogeneous mixtures was increased from lower concentrations (5.88–30%) to higher concentrations (≥80%), the colour of the mixture changed from dark brown to brownish yellow. This colour change may indicate that the structure of iron in mixtures at x(H₂O) ≤ 30% may be different from that at x(H₂O) ≥ 80%, because the colour of each solution came from an optical frequency characteristic of a specific iron structure. Hence, the formation of homogeneous [bmim]Fe(III)Cl₄–H₂O mixtures possibly affects the oxidative absorption of hydrogen sulfide.

2.2. Acidity of Homogeneous [bmim]Fe(III)Cl₄–H₂O Mixtures

Solution pH is widely recognized as a key variable for H₂S oxidation process using iron chelates [20]. Acidity is also a key variable for oxidation of H₂S by [bmim]Fe(III)Cl₄ [21]. Pyridine has been used as a probe molecule for the determination of the Lewis and Brønsted acidities of solid acids and ionic liquids by monitoring the bands in the range of 1400–1700 cm⁻¹ arising from its ring vibration modes [22]. The principle of this experiment is as follows: if pyridine, a weak Lewis base, is mixed with acid, the interaction between these two compounds will exhibit a correlated band in IR spectra [23]. In this work, pyridine was added to homogeneous [bmim]Fe(III)Cl₄–H₂O mixtures, then their acidity was measured by infrared spectroscopy.

As shown in Figure 2, when x(H₂O) varies from 0 (pure [bmim]Fe(III)Cl₄) to 30%, the IR bands at around 1540 and 1635 cm⁻¹ are attributed to Brønsted acids [22], while the IR bands at around 1487 and 1610 cm⁻¹ are attributed to Lewis acids [21,24]. These assignments demonstrate that the homogenous mixtures at x(H₂O) ≤ 30% are simultaneously Brønsted and Lewis acids. In contrast, when x(H₂O) reaches 80%, the intensity of the IR bands at 1540 and 1635 cm⁻¹ decrease, and are almost absent at x(H₂O) = 90%. However, the IR bands at 1487 and 1610 cm⁻¹ are still present. These observations indicate that increasing the content of water causes a decrease in the amount of Brønsted acid present, resulting in the presence of almost exclusively Lewis acids at x(H₂O) ≥ 80%. Since the H at the 2 position on the imidazolium ring in [bmim]Fe(III)Cl₄ is easily removed in Brønsted acids [25], the absence of Brønsted acids was mainly caused by the decomposition of [bmim]Fe(III)Cl₄, as discussed below. On the other hand, Lewis acids resulted from FeCl₃ [23]. Therefore, it is possible that the FeCl₃ concentration decreased as the water content increased, causing a decrease in the quantity of Lewis acids present. For example, when x(H₂O) increases from 80% to 90%, the pH of the homogeneous mixtures increases from 1.17 to 1.28. However, the acidity of the mixtures at x(H₂O) ≤ 30% is too low (negative pH) to be precisely measured by a pH Meter.

2.3. Raman Spectral Analysis of Homogeneous [bmim]Fe(III)Cl₄–H₂O Mixtures

Raman spectra were used to investigate the structure of iron species in the homogeneous [bmim]Fe(III)Cl₄–H₂O mixtures. The Raman spectra for the homogeneous mixture at x(H₂O) = 5.88% are shown in Figure 3. The peaks at 110, 136, 333, and 390 cm⁻¹ correspond very well with literature values for Fe(III)Cl₄⁻ in concentrated aqueous solution [26], indicating that Fe(III)Cl₄⁻ is the predominant iron-containing species in the homogeneous mixtures at x(H₂O) = 5.88% before oxidative absorption of H₂S. Since homogeneous liquids at x(H₂O) ≤ 30% are all dark brown, which corresponds to the colour of [bmim]Fe(III)Cl₄ [1], [bmim]Fe(III)Cl₄ was the dominating species at x(H₂O) ≤ 30%.

Unfortunately, the peaks at 136, 333, and 390cm⁻¹ disappeared for the homogeneous mixture containing 80% H₂O. A new peak at 318 cm⁻¹ appeared (see Figure 3), belonging to FeCl₃·2H₂O [26]. This peak provides strong evidence for the decomposition of [bmim]Fe(III)Cl₄ into a new species, such as FeCl₃·2H₂O (see Equation(2)) at x(H₂O) = 80%, which is consistent with Liu’s research [27].

(2)

The colour of the homogeneous mixture at x(H₂O) ≥ 80% demonstrates that FeCl₃·2H₂O is the iron compound present at x(H₂O) ≥ 80%. Since both FeCl₃·2H₂O and hydrophilic [bmim]Cl [17] are soluble in water, the homogeneous mixtures at x(H₂O) ≥ 80% are true solutions. These results come in complete opposition with previous works asserting that due to the hydrophobic property of [bmim]Fe(III)Cl₄, the homogeneous mixture of [bmim]Fe(III)Cl₄ and H₂O may be an emulsion rather than a true solution [12].

2.4. Oxidative Absorption of H₂S by Homogeneous [bmim]Fe(III)Cl₄–H₂O Mixtures

Oxidative absorption of H₂S by homogeneous [bmim]Fe(III)Cl₄–H₂O mixtures was investigated at 60 °C. Figure 3 and Figure 4 provide a good picture of the oxidative product. As can be seen in Figure 4, the characteristic feature of FeCl₃·2H₂O at 318 cm⁻¹ in homogenous mixtures of 80% H₂O disappears after oxidative absorption of H₂S. Simultaneously, a new peak at 153 cm⁻¹ belonging to FeCl₂ [28], as well as new peaks at 219 and 473 cm⁻¹, both due to S₈ [29], appears. These results indicate that FeCl₃·2H₂O was completely reduced to FeCl₂. Therefore, the reaction can be summarized by Equation (3).

16FeCl₃·2H₂O + 8H₂S → 16FeCl₂ + S₈ + 16HCl + 32H₂O

(3)

Equation (3) shows that HCl is generated from the oxidative absorption of H₂S by homogeneous mixtures at x(H₂O) ≥ 80%. Consequently, it is apparent that the acidity of mixtures at x(H₂O) ≥ 80% is improved after oxidative absorption of H₂S and discharge of volatile HCl from the acidic mixtures. Furthermore, as shown in Figure 5c,d, homogeneous mixtures at x(H₂O) = 90% after oxidative-absorption of H₂S become colourless due to the complete reduction of FeCl₃·2H₂O and S₈, the oxidative product of H₂S, is suspended in the mixtures. Figure 6 shows that the surface tension of homogeneous mixtures at x(H₂O) ≤ 30% remains almost constant up to about 37 mN·m⁻¹, whereas, a remarkable increase in the surface tension is observed for homogeneous mixtures at x(H₂O) = 80%, moreover, for homogeneous mixtures at x(H₂O) = 90%, the surface tension sharply increase to 56.57 mN·m⁻¹. At x(H₂O) ≥ 80%, the surface tension of homogeneous mixtures increases remarkably, resulting in the suspension of S₈ in the mixtures.

However, it is difficult to prevent the suspended S₈ from the formation of sulfur blockage. Therefore, catalytic-oxidation of H₂S by homogeneous mixtures at x(H₂O) ≥ 80% suffers from at least two major drawbacks: (1) the slow rate of oxidation of free Fe(II) ions, such as FeCl₂ oxidation by O₂ [30], and (2) the sulfur blockage of suspended S₈.

On the contrary, with the color deepening of homogeneous mixtures at x(H₂O) = 30% after oxidative absorption of H₂S, stable deposits of elemental sulfur, which is easy to be separated, is likely produced, as shown in Figure 5b. However, the characteristic peaks at 110, 136, 333, and 390 cm⁻¹ for FeCl₄⁻ also exist in the homogeneous mixture at x(H₂O) = 5.88% after the absorption of H₂S with the appearance of no new peaks in the IR spectra (see Figure 3). These results indicate that Fe(III)Cl₄⁻ is also an essential component in the homogeneous mixture after oxidative absorption of H₂S. In fact, Fe(III)Cl₄⁻ was probably changed to Fe(II)Cl₄²⁻ [6], and H₂S was oxidized to S₈ [24]. But the concentration of Fe(II)Cl₄²⁻ and sulfur in mixtures at x(H₂O) = 5.88% after oxidative-absorption of H₂S were too low to be detected by Raman spectroscopy (see

Table 1. Effect of water concentration on the weight percent of absorbed H₂S and volatile HCl emissions in homogenous [bmim]Fe(III)Cl₄–H₂O mixtures at 60 °C.

Water Concentration in Homogeneous Mixtures (v/v) %	0	5.88	80
Weight percent of iron(III) in homogeneous mixtures (%)	16.6	15.87	4.19
Weight percent of absorbed H2S in mixtures (%)	0.78	0.82	1.15
Concentration of volatile HCl in mixtures (×10−6)	841.11	326.86	52.16

). This was strengthened by the previous studies [21]. The Raman characteristic peaks of S₈ and Fe(II)Cl₄²⁻ after oxidative-absorption of H₂S by pure [bmim]FeCl₄ were not observed. However, at lower concentration of [bmim]FeCl₄ in the mixtures of [bmim]Fe(III)Cl₄, [bmim]Cl and water, the acidic strength decreased remarkably, resulting in higher concentration of Fe(II)Cl₄²⁻ and sulfur, and hence the new features of S₈ and Fe(II)Cl₄²⁻ after oxidative-absorption of H₂S by the mixtures of [bmim]Fe(III)Cl₄, [bmim]Cl and water could be detected by a Raman Spectrometer.

In order to get the more catalytic oxidation product of H₂S by homogeneous mixtures, the first absorption of H₂S for 20 min and the second regeneration of O₂ for 40 min in the homogeneous mixtures at x(H₂O) = 5.88% were conducted for six consecutive cycles, then, the homogeneous mixtures at x(H₂O) = 5.88% were filtered and vacuum-dried. We get an amount of yellow powder. The XRD patterns of the yellow powder are shown in Figure 7. The yellow powder is observed with significant S (23) peak reflection demonstrates that the powder is elemental sulfur.

Our conclusion was strengthened by the fact that the product of oxidative absorption of H₂S using pure [bmim]Fe(III)Cl₄ was sulfur [24].

In addition, Fe(III)Cl₄⁻ in DMSO was reduced to Fe(II)Cl₄²⁻ in the oxidation of H₂S [31]. Finally, our previous research also showed H₂S was converted to S₈ by [bmim]FeCl₄ of [bmim]Fe(III)Cl₄–[bmim]Cl–H₂O mixtures, and [bmim]FeCl₄ in the process was converted to [bmim]Fe(II)Cl₄H [21]. Hence, the oxidation of H₂S by homogeneous mixtures at x(H₂O) ≤ 30% is written as Equation (4).

(4)

Nguyen et al. suggested that [bmim]FeCl₄ could be reduced to give ([bmim]FeCl₃)*, which could be transformed into [Fe(II)Cl]⁺[(bmim)Cl₂]⁻ [6]. Similarly, our previous research also suggested [bmim]Fe(III)Cl₄ was firstly reduced by H₂S to give ([bmim]FeCl₃)*, which was then transformed into [bmim]Fe(II)Cl₄H as soon as it reacted with HCl (see Equations (5) and (6)).

16[bmim]Fe(III)Cl₄ + 8H₂S → 16([bmim]Fe(II)Cl₃)* + S₈ + 16HCl

(5)

([bmim]Fe(II)Cl₃)* + HCl → [bmim]Fe(II)Cl₄H

(6)

From Equation (4), an increase in the acidity of homogeneous mixtures after oxidation of H₂S is apparent due to the production of H⁺. In addition, Fe(III)Cl₄⁻ has a longer Fe(III)–Cl bond distance than FeCl₃ [6]. Therefore, it is likely that after oxidation of H₂S, H⁺ and Cl⁻ present in homogeneous mixtures becomes HCl. The existence of HCl is also strengthened by Equation (5), which shows that HCl comes from the intermediate species.

Trace amounts of volatile HCl from the mixtures during oxidative absorption of H₂S were indeed detected by ion chromatography. Effect of water concentration on oxidative absorption of H₂S by [bmim]Fe(III)Cl₄–H₂O mixtures without other additives at 60 °C was shown in Table 1. Table 1 shows an increase in weight percent of absorbed H₂S and a decrease in volatile HCl emissions with increasing water concentration were seen.

Previous studies demonstrated that the oxidation of H₂S by Fe³⁺ or an iron(III) chelate is a fast chemical reaction [20,32]. As soon as H₂S was absorbed, it was oxidized. In addition, many researchers have observed a high solubility of H₂S in a series of 1-butyl-3-methylimidazolium ([bmim]⁺)-based ionic liquids due to the formation of H₂S–Cl⁻ or other complexes [33,34,35]. It can be seen from Table 1 that the concentration of iron(III) in mixtures is very high in comparison with chelated iron solutions [36]. A high concentration of absorbed H₂S in these mixtures is not seen, however. On the contrary, the concentration of iron decreases as the water content increases, resulting in an increase of the weight percent of absorbed H₂S. A key factor is that the [bmim]Fe(III)Cl₄–H₂O mixtures were not only an oxidizer but also a solution, thus, the low concentrations of absorbed H₂S in mixtures was most likely caused by the strong acidity of the mixtures. Figure 2 shows that the homogenous mixtures at x(H₂O) ≤ 30% are simultaneously Brønsted and Lewis acids. Moreover, their acidity was too strong to be measured by a pH Meter. Even when the water content of the mixtures increases to 80%, the pH only reaches 1.17. However, a highly acidic environment was disadvantageous to the solubility of H₂S [20,32]. Consequently, the weight percent of absorbed H₂S in mixtures was very low in comparison with iron concentrations. Nevertheless, the decrease in acidity by increasing the water content improved the solubility of H₂S in the mixtures, and the weight percent of absorbed H₂S in mixtures was improved.

Similarly, decreasing acidity by increasing water content enhances the solubility of acidic HCl, and, thus, volatile HCl emissions decreased with increasing water content. However, during oxidative absorption of H₂S by homogenous [bmim]Fe(III)Cl₄–H₂O mixtures at 60 °C increasing water also led to more evaporation of water, which decreased the weight percent of absorbed H₂S in mixtures, according to the absorption measurements. Volatile HCl emissions and the evaporation of water indicated that the actual weight percent of absorbed H₂S in mixtures was higher than the measured value.

It’s worth mentioning that increasing x(H₂O) was favourable for increasing oxidative-absorption of H₂S and decreasing volatile HCl emissions, but suspended S₈ generated at x(H₂O) ≥ 80% was difficult to separate from the mixtures and the formation of the sulfur blockage.

3. Materials and Methods

3.1. Preparation of [bmim]Fe(III)Cl₄–H₂O Mixtures

[bmim]Fe(III)Cl₄ was synthesized via a literature method [1]. Mixtures of [bmim]Fe(III)Cl₄ and water were prepared by adding water to [bmim]Fe(III)Cl₄ with vigorous stirring for 24 h. Ten [bmim]Fe(III)Cl₄–H₂O mixtures were prepared in volume fractions of x(H₂O), where x(H₂O) = 5.88%, 10%, 20%, 30%, 40%, 50%, 60%, 70%, 80%, and 90%.

3.2. Instruments

The acidity of each mixture was measured by a TENSOR 27 FT-IR Spectrometer (Bruker Optics, Ettlingen, Germany) using a method of determining the acidity of [bmim]Fe(III)Cl₄ ionic liquids [24]. Pyridine was used as a base infrared spectroscopy probe molecule to determine the acidity of ionic liquids. All of the infrared spectroscopy samples were prepared by mixing pyridine and ionic liquids at the ratio of one to five and then spreading as liquid films on KBr windows.

The iron structure of each mixtures was obtained using a laser confocal spectrometer (model labRAM Aramis, Horiba Jobin Yivon, France) with a He–Ne laser (632.8nm)). Raman samples were measured in glass capillaries with an inner diameter of 0.9–1.1 mm and 0.10–0.15 mm wall thickness.

The surface tension of each mixtures was measured by OCA 15 (Dataphysics, Filderstadt, Germany) at room temperature.

XRD of the powder was obtained using an X ray diffractometer (model Ultima IV, Rigaku, Tokyo, Japan).

3.3. Absorption Measurements

The apparatus of oxidation-absorption of H₂S in ionic liquids has been presented in detail [21]. H₂S from a commercial gas cylinder was first bubbled at a pressure of 0.05 Mpa and a flow rate of 15 mL/min through predetermined amounts of each liquid (about 3.5–4 g) loaded in a glass bubbling absorption tube (140 mm long, 20 mm maximum inner diameter), which was placed in a water bath at 60 °C, then bubbled into 200 mL of ultra-pure water in a tail gas absorber.

After 20 min, the glass bubbling absorption tube was weighed using an AL104 precision balance (Mettler Toledo, Shanghai, China) with an uncertainty of ±0.0001 g, to measure the mass of H₂S absorbed. From this mass, the H₂S absorbed per gram of ionic liquid was calculated. Then, a DX-600 ion chromatography apparatus (DIONEX Co., Sunnyvale, CA, USA) was used to analyse the Cl⁻ concentration in ultra-pure water, which had been absorbed along with H₂S. From this data, the concentration of volatile HCl per gram of ionic liquid was calculated.

If the ionic liquids absorbed H₂S was regenerated by O₂, O₂ was bubbled at a pressure of 0.05 Mpa and a flow rate of 30 mL/min through the ionic liquids absorbed H₂S for 40 min.

4. Conclusions

We have prepared homogeneous mixtures of [bmim]Fe(III)Cl₄–H₂O at x(H₂O) = 5.88–30% and above 80%, which showed the presence of both Brønsted and Lewis acids at x(H₂O) ≤ 30%, and only Lewis acids at x(H₂O) ≥ 80% by FT-IR. Raman spectra showed Fe(III)Cl₄⁻ was the dominating iron species at x(H₂O) ≤ 30%, which was reduced by H₂S to Fe(II)Cl₄⁻. FeCl₃·2H₂O was the primary iron compound at x(H₂O) ≥ 80%, which was reduced to FeCl₂. In addition, oxidative-absorption of H₂S by homogeneous [bmim]Fe(III)Cl₄–H₂O mixtures demonstrated that increasing x(H₂O) increased the weight percent of absorbed H₂S and decreased the volatile HCl emissions. However, there are some disadvantages in oxidation of H₂S by homogeneous [bmim]Fe(III)Cl₄–H₂O mixtures at x(H₂O) ≥ 80%, including sulfur blockage of suspended S₈ and the slow regenerative rate of FeCl₂. Therefore, it is necessary that controlling the x(H₂O) below 80% ensures that oxidative-absorption of H₂S by [bmim]Fe(III)Cl₄ is performed well.