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

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.

[bmim]Fe(III)Cl 4 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)Cl 4 as an alternative to an iron(III) chelating solution for catalytic oxidation of H 2 S. 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)Cl 4 , however, do not require a complexing agent. In addition, sulfur oxo-acid salts cannot form in acidic media [9]. Catalytic-oxidation of H 2 S by [bmim]Fe(III)Cl 4 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].
Apart from sulfur, which is easily recovered by filtration, only H 2 O is generated. Since [bmim]Fe(III)Cl 4 is hydrophobic [1], mixtures of [bmim]Fe(III)Cl 4 and water should form two phases that should be easily separable. However, Lee et al. reported that 20% (v/v) of [bmim]Fe(III)Cl 4 and water were fully miscible after vigorous shaking [11]. Wang [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)Cl 4 . 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)Cl 4 has an influence on the oxidative absorption of H 2 S.
In this paper, we report our findings on the properties of homogeneous mixtures of [bmim]Fe(III)Cl 4 and H 2 O, and their effect on the oxidative absorption of H 2 S. Additional work on the regeneration of [bmim]Fe(III)Cl 4 from these aqueous homogeneous mixtures will be discussed in a future paper.

Effect of Water Concentration on the Formation of [bmim]Fe(III)Cl 4 -H 2 O Mixtures
The effect of water concentration on the formation of homogeneous [bmim]Fe(III)Cl 4 -H 2 O mixtures was studied by adding different amounts of water to [bmim]Fe(III)Cl 4 . Figure 1 shows that dark brown homogeneous liquids are obtained at x(H 2 O) values from 5.88% to 30%. At x(H 2 O) 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(H 2 O) ≥ 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)Cl 4 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)Cl 4 -H 2 O mixtures without evaporation method [11,12]  x(H2O) from left to right are: 5.88%, 10%, 20%, 30%, 40%, 50%, 60%, 70%, 80%, and 90%.

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.

Acidity of Homogeneous [bmim]Fe(III)Cl 4 -H 2 O Mixtures
Solution pH is widely recognized as a key variable for H 2 S oxidation process using iron chelates [20]. Acidity is also a key variable for oxidation of H 2 S by [bmim]Fe(III)Cl 4 [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)Cl 4 -H 2 O mixtures, then their acidity was measured by infrared spectroscopy.
As shown in Figure 2, when x(H 2 O) varies from 0 (pure [bmim]Fe(III)Cl 4 ) 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(H 2 O) ≤ 30% are simultaneously Brønsted and Lewis acids. In contrast, when x(H 2 O) reaches 80%, the intensity of the IR bands at 1540 and 1635 cm −1 decrease, and are almost absent at x(H 2 O) = 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(H 2 O) ≥ 80%.
Since the H at the 2 position on the imidazolium ring in [bmim]Fe(III)Cl 4 is easily removed in Brønsted acids [25], the absence of Brønsted acids was mainly caused by the decomposition of [bmim]Fe(III)Cl 4 , as discussed below. On the other hand, Lewis acids resulted from FeCl 3 [23]. Therefore, it is possible that the FeCl 3 concentration decreased as the water content increased, causing a decrease in the quantity of Lewis acids present. For example, when x(H 2 O) 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(H 2 O) ≤ 30% is too low (negative pH) to be precisely measured by a pH Meter. x(H2O) from left to right are: 5.88%, 10%, 20%, 30%, 40%, 50%, 60%, 70%, 80%, and 90%.

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.

Raman Spectral Analysis of Homogeneous [bmim]Fe(III)Cl 4 -H 2 O Mixtures
Raman spectra were used to investigate the structure of iron species in the homogeneous [bmim]Fe(III)Cl 4 -H 2 O mixtures. The Raman spectra for the homogeneous mixture at x(H 2 O) = 5.88% are shown in Figure 3.
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].

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. Figures 3 and 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). Unfortunately, the peaks at 136, 333, and 390cm −1 disappeared for the homogeneous mixture containing 80% H 2 O. A new peak at 318 cm −1 appeared (see Figure 3), belonging to FeCl 3 ·2H 2 O [26]. This peak provides strong evidence for the decomposition of [bmim]Fe(III)Cl 4 into a new species, such as FeCl 3 ·2H 2 O (see Equation (2)) at x(H 2 O) = 80%, which is consistent with Liu's research [27].
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].

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. Figures 3 and 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).  Figures 3 and 4 provide a good picture of the oxidative product. As can be seen in Figure 4, the characteristic feature of FeCl 3 ·2H 2 O at 318 cm −1 in homogenous mixtures of 80% H 2 O disappears after oxidative absorption of H 2 S. Simultaneously, a new peak at 153 cm −1 belonging to FeCl 2 [28], as well as new peaks at 219 and 473 cm −1 , both due to S 8 [29], appears. These results indicate that FeCl 3 ·2H 2 O was completely reduced to FeCl 2 . Therefore, the reaction can be summarized by Equation (3).
Equation (3) shows that HCl is generated from the oxidative absorption of H 2 S by homogeneous mixtures at x(H 2 O) ≥ 80%. Consequently, it is apparent that the acidity of mixtures at x(H 2 O) ≥ 80% is improved after oxidative absorption of H 2 S and discharge of volatile HCl from the acidic mixtures. Furthermore, as shown in Figure 5c,d, homogeneous mixtures at x(H 2 O) = 90% after oxidative-absorption of H 2 S become colourless due to the complete reduction of FeCl 3 ·2H 2 O and S 8 , the oxidative product of H 2 S, is suspended in the mixtures. Figure 6 shows that the surface tension of homogeneous mixtures at x(H 2 O) ≤ 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 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.  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. However, it is difficult to prevent the suspended S 8 from the formation of sulfur blockage. Therefore, catalytic-oxidation of H 2 S by homogeneous mixtures at x(H 2 O) ≥ 80% suffers from at least two major drawbacks: (1) the slow rate of oxidation of free Fe(II) ions, such as FeCl 2 oxidation by O 2 [30], and (2) the sulfur blockage of suspended S 8 .  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 5(b). 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)Cl4 2− [6], and H2S was oxidized to S8 [24]. But the concentration of Fe(II)Cl4 2− 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). This was strengthened by the previous studies [21]. 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.  On the contrary, with the color deepening of homogeneous mixtures at x(H 2 O) = 30% after oxidative absorption of H 2 S, 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 FeCl 4 − also exist in the homogeneous mixture at x(H 2 O) = 5.88% after the absorption of H 2 S with the appearance of no new peaks in the IR spectra (see Figure 3). These results indicate that Fe(III)Cl 4 − is also an essential component in the homogeneous mixture after oxidative absorption of H 2 S. In fact, Fe(III)Cl 4 − was probably changed to Fe(II)Cl 4 2− [6], and H 2 S was oxidized to S 8 [24].
But the concentration of Fe(II)Cl 4 2− and sulfur in mixtures at x(H 2 O) = 5.88% after oxidative-absorption of H 2 S were too low to be detected by Raman spectroscopy (see Table 1). This was strengthened by the previous studies [21]. The Raman characteristic peaks of S 8  In order to get the more catalytic oxidation product of H 2 S by homogeneous mixtures, the first absorption of H 2 S for 20 min and the second regeneration of O 2 for 40 min in the homogeneous mixtures at x(H 2 O) = 5.88% were conducted for six consecutive cycles, then, the homogeneous mixtures at x(H 2 O) = 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.  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 5(b). 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)Cl4 2− [6], and H2S was oxidized to S8 [24]. But the concentration of Fe(II)Cl4 2− 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). This was strengthened by the previous studies [21]. 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 H 2 S using pure [bmim]Fe(III)Cl 4 was sulfur [24].
In 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)Cl4 2− 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). 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 Fe 3+ 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.   (5) and (6) From Equation (4), an increase in the acidity of homogeneous mixtures after oxidation of H 2 S is apparent due to the production of H + . In addition, Fe(III)Cl 4 − has a longer Fe(III)-Cl bond distance than FeCl 3 [6]. Therefore, it is likely that after oxidation of H 2 S, 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 H 2 S were indeed detected by ion chromatography. Effect of water concentration on oxidative absorption of H 2 S by [bmim]Fe(III)Cl 4 -H 2 O mixtures without other additives at 60 • C was shown in Table 1. Table 1 shows an increase in weight percent of absorbed H 2 S and a decrease in volatile HCl emissions with increasing water concentration were seen. Previous studies demonstrated that the oxidation of H 2 S by Fe 3+ or an iron(III) chelate is a fast chemical reaction [20,32]. As soon as H 2 S was absorbed, it was oxidized. In addition, many researchers have observed a high solubility of H 2 S in a series of 1-butyl-3-methylimidazolium ([bmim] + )-based ionic liquids due to the formation of H 2 S-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 H 2 S 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 H 2 S. A key factor is that the [bmim]Fe(III)Cl 4 -H 2 O mixtures were not only an oxidizer but also a solution, thus, the low concentrations of absorbed H 2 S in mixtures was most likely caused by the strong acidity of the mixtures. Figure 2 shows that the homogenous mixtures at x(H 2 O) ≤ 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 H 2 S [20,32]. Consequently, the weight percent of absorbed H 2 S in mixtures was very low in comparison with iron concentrations. Nevertheless, the decrease in acidity by increasing the water content improved the solubility of H 2 S in the mixtures, and the weight percent of absorbed H 2 S 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 H 2 S by homogenous [bmim]Fe(III)Cl 4 -H 2 O mixtures at 60 • C increasing water also led to more evaporation of water, which decreased the weight percent of absorbed H 2 S in mixtures, according to the absorption measurements. Volatile HCl emissions and the evaporation of water indicated that the actual weight percent of absorbed H 2 S in mixtures was higher than the measured value.
It's worth mentioning that increasing x(H 2 O) was favourable for increasing oxidative-absorption of H 2 S and decreasing volatile HCl emissions, but suspended S 8 generated at x(H 2 O) ≥ 80% was difficult to separate from the mixtures and the formation of the sulfur blockage.

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)Cl 4 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).

Absorption Measurements
The apparatus of oxidation-absorption of H 2 S in ionic liquids has been presented in detail [21]. H 2 S 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 H 2 S absorbed. From this mass, the H 2 S 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 H 2 S. From this data, the concentration of volatile HCl per gram of ionic liquid was calculated.
If the ionic liquids absorbed H 2 S was regenerated by O 2 , O 2 was bubbled at a pressure of 0.05 Mpa and a flow rate of 30 mL/min through the ionic liquids absorbed H 2 S for 40 min.

Conclusions
We