Performance and Kinetics of BPA Degradation Initiated by Powdered Iron (or Ferrous Sulfate) and Persulfate in Aqueous Solutions

Abstract

The widespread use of bisphenol A (BPA) in industry has resulted in BPA contamination of water bodies and even endocrine-disrupting effects on organisms and humans through water transmission. Advanced oxidation processes based on sulfate radicals have received increasing attention due to their ability to efficiently degrade endocrine disruptors (including BPA) in water. In this study, powdered iron (Fe(0)) and ferrous sulfate (Fe(II)) were used as activators to activate persulfate (PS) for the degradation of BPA. The effects of the dosage of the activator, the concentration of PS, the concentration of BPA, the initial solution pH, and the reaction temperature on the degradation efficiency of BPA in Fe(II)/PS and Fe(0)/PS systems were investigated, and the kinetics of BPA degradation under different reaction conditions were analyzed. The results showed that the optimal conditions were [Fe(II)] = 0.1 g/L, [PS] = 0.4 mM, [BPA] = 1 mg/L, T = 70 °C and pH = 5.0 for the Fe(II)/PS system and [Fe(0)] = 0.5 g/L, [PS] = 0.5 mM, [BPA] = 1 mg/L, T = 70 °C and pH = 5.0 for the Fe(0)/PS system; both systems were able to achieve equally good degradation of BPA. The degradation of BPA in the Fe(II)/PS system satisfied the pseudo-secondary kinetic equation under varying PS concentration conditions, otherwise the degradation of BPA in both systems conformed to the pseudo-first-order kinetic equation.

1. Introduction

Due to the rapid development of the chemical industry, many new chemical substances are constantly being synthesized for industrial, agricultural, and pharmaceutical use. Although these substances have brought benefits to the public, they have also caused serious environmental pollution issues. Bisphenol A (BPA) is an important organic chemical raw material used in the preparation of epoxy resins and polycarbonates which are widely used in production and life [1]. Many products contain BPA, such as plastic containers, thermal paper, coatings on cans, disposable paper cups, etc. Trace amounts of BPA are frequently detected in surface water due to the discharge of industrial wastewater or domestic sewage [2]. It is now clear that BPA is an environmental hormone that has a disruptive effect on the endocrine system of living organisms [3,4,5]. It is essential to reduce the concentration of BPA in drinking water sources to ensure the safety of drinking water. As the conventional treatment technologies (adsorption, filtration, and biochemical oxidation) are not completely effective in removing BPA [6,7,8,9], this poses a challenge to the treatment of BPA via advanced methods.

The removal of recalcitrant organic matter (including endocrine disruptor BPA) by advanced oxidation processes (AOPs) has become one of the hot spots in water treatment research. The AOPs mainly use highly oxidizing free radicals generated in situ to degrade organic matters, such as hydroxyl radicals (•OH, E⁰ = 1.9−2.7) and sulfate radicals (SO₄^•−, E⁰ = 2.5−3.1 V) [10,11]. Compared with hydroxyl radicals, sulfate radicals have a higher oxidation capacity for organic substances and are more adaptable to reaction conditions. Therefore, sulfate radical oxidation is more advantageous than the conventional hydroxyl radical oxidation in the treatment of recalcitrant organics [12]. Oxidants (hydrogen peroxide [13], ozone [14], persulfate (PS) [15], and peroxymonosulfate [16]) can be used as precursors to generate free radicals with the help of all kinds of energies or materials [17]. The means of activation can be ultraviolet light (UV) [18], heat [19], microwaves (MW) [20], ultrasound (US) [21], heavy metals (Fe [22], Co [23], Cu [24], Mn [25]), and carbon materials [26,27]. Among the four oxidants mentioned above, PS is favored by researchers because of its stable nature and ease of transport and storage [28]. PS alone (S₂O₈²⁻, E⁰ = 2.01 V) has a weak ability to directly oxidize organic matters, while it can be activated to produce free radicals to acquire better performance toward organics degradation [29]. The molecule of PS contains an O—O bond (140 kJ/mol, 1.497 Å), which can be broken to form sulfate radicals at external energies (UV, US, MW, and heat) above the O—O bond energy [30]. In addition, PS can also undergo redox reaction with transition metals to form sulfate radicals [31]. As iron is a common transition metal that is abundant on earth, inexpensive, and non-toxic, it is considered as a benign catalyst used in AOPs [32]. For example, divalent iron (Fe(II)) could catalyze PS to generate sulfate radicals [33]. In addition, powdered iron (Fe(0)) has also been used to activate PS in removing organic matter [34].

Although the use of Fe(II) and Fe(0) for activating PS is relatively common in research, the economic and environmental friendliness of iron is very prominent. Iron-based PS systems have an advantage over other transition metals-based PS systems (Co/PS, Cu/PS, Mn/PS) in practical applications toward wastewater treatment. Fewer studies have been conducted using Fe(II)/PS and Fe(0)/PS systems to treat endocrine disruptor BPA; therefore, a comparative study of the two systems can provide favorable parameters for future applications of iron-based systems in actual wastewater treatment. In this study, Fe(0) and Fe(II) were used to remove BPA from an aqueous solution in the presence of PS. The abilities of two systems to remove BPA under different influencing factors were investigated and the corresponding reaction kinetic models were developed. At the same time, a profound analysis of the mechanism was carried out in terms of catalysis and degradation in Fe(0)/PS and Fe(II)/PS systems.

2. Results

2.1. Dosage of Activators

Figure 1a,b shows the effects of different dosages of Fe(II) and Fe(0) on activating PS in debating BPA, respectively. It can be seen in Figure 1a that PS was able to oxidize 19.86% of BPA without the addition of catalyst Fe(II), which meant that PS alone could oxidize organic matter at a slower rate. The addition of 0.1 g/L of Fe(II) was able to achieve an optimal BPA removal (C/C₀ = 0.455) in the presence of PS. Subsequently, the dosage of Fe(II) was increased from 0.2 to 0.5 g/L, which brought about a gradual decrease in the BPA removal rate (0.5 g/L of Fe(II) corresponding to C/C₀ = 0.706). This indicated that too much Fe(II) did not favor the catalytic reaction rate. The reaction of Fe(II) activating PS to generate SO₄^•− is shown in Equation (1) [35]. From this equation, Fe(II) is able to be oxidized to Fe(III) by PS, while PS is reduced to SO₄^•− and SO₄²⁻. However, the excess Fe(II) will react with the SO₄^•− formed in the water, converting SO₄^•− to SO₄²⁻ (Equation (2)) [35,36] and thus losing the ability to degrade organic matter. Therefore, the optimization of the Fe(II) and PS dosing ratio is beneficial to the improvement of degradation efficiency. Figure 1b exhibits the trend of BPA degradation in Fe(0)/PS system. In contrast to the trend in Figure 1a, the residual concentration of BPA decreased gradually as the dosage of Fe(0) was increased from 0 to 0.5 g/L. The best BPA removal (C/C₀ = 0.514) was achieved within 60 min with 0.5 g/L of Fe(0). While for the Fe(II)/PS system, the best BPA removal was acquired at Fe(II) of 0.1 g/L (C/C0 = 0.455). Since the activation of Fe(0) on PS is heterophasic, whereas the activation of Fe(II) on PS is homophasic, the performance on the optimal concentrations of the two catalysts is not consistent. In fact, the activation of PS by Fe(0) is an indirect process. Fe(0) should first be converted to Fe(II) before the reaction of Equation (1) occurs to activate PS. Fe(0) can be converted to Fe(II) under three conditions, see Equations (3)–(5) [37]. Since Fe(0) catalyzes PS in a heterogeneous manner, Fe(0) does not dissolve into solution quickly and is not easily overdosed compared to Fe(II), which becomes an advantage of Fe(0) in activating PS. Therefore, Fe(0) is able to avoid the problem of quenching SO₄^•− by generating too much Fe(II) in the reaction.

$${\mathrm{Fe}}^{2+}+{\mathrm{S}}_2{\mathrm{O}}_8^{2-}\to {\mathrm{Fe}}^{3+}+{\mathrm{SO}}_4^{2-}+{\mathrm{SO}}_4^{•-}$$

(1)

$${\mathrm{Fe}}^{2+}+{\mathrm{SO}}_4^{•-}\to {\mathrm{Fe}}^{3+}+{\mathrm{SO}}_4^{2-}$$

(2)

$${\mathrm{Fe}}^0+{\mathrm{S}}_2{\mathrm{O}}_8^{2-}\to {\mathrm{Fe}}^{2+}+{\mathrm{SO}}_4^{2-}$$

(3)

$${\mathrm{Fe}}^0+{\mathrm{H}}_2\mathrm{O}\to {\mathrm{Fe}}^{2+}+{2\mathrm{OH}}^{-}+{\mathrm{H}}_2$$

(4)

$${\mathrm{Fe}}^0+{\mathrm{H}}_2\mathrm{O}+{0.5\mathrm{O}}_2\to {\mathrm{Fe}}^{2+}+{2\mathrm{OH}}^{-}$$

(5)

Figure 2a,b develops the kinetic models of BPA degradation by Fe(II)/PS and Fe(0)/PS, respectively. The degradation of BPA at varying dosages of Fe(II) and Fe(0) was in accordance with the pseudo first-order kinetic equations. The correlation coefficients (R²) of the established equations were all above 0.93. In Figure 2a, when Fe(II) was 0.1 g/L, the degradation of BPA obtained the maximum rate of 0.01171 min⁻¹. In Figure 2b, the maximum BPA degradation rate occurred at 0.5 g/L of Fe(0).

2.2. Concentrations of PS

Figure 3a depicts the changes in BPA removal at different PS concentrations with the addition of Fe(II). As can be seen from Figure 3a, Fe(II) had no ability to remove BPA at all when the PS concentration was 0 mM. Unlike Fe(II), Fe(0) exhibited limited removal of BPA (60 min, C/C₀ = 0.801) in the absence of PS (Figure 3b), which should be attributed to the adsorption and reduction of BPA by nano-sized Fe(0) possessing a large specific surface area. In Figure 3a, as the PS concentration increased from 0.2 mM to 0.4 mM, the residual concentration of BPA in the solution continued to decrease. However, at a PS concentration of 0.5 mM, it was found that the change in BPA removal was not significant compared to that at a PS concentration of 0.4 mM. This indicated that the best treatment effect of BPA was achieved at PS of 0.4 mM. Similar results were obtained in a previous study which showed that a high concentration PS could inhibit the degradation of organics in Fe(II)/PS system [38]. Although PS as an oxidant could provide SO₄^•− through Fe(II) catalyzing, excess PS could react with the SO₄^•− (Equations (6)) [39,40] and reduce the amount of SO₄^•− in the solution, thus inhibiting the degradation of organics. However, it was found from Figure 3b that the increase in Fe(0) concentration from 0.1 mM to 0.5 mM did not bring about the inhibition of BPA degradation, and the best BPA removal was achieved at 0.5 mM of PS. Compared to Fe(II) for PS catalysis, 0.5 mM of PS was not excessive in the Fe(0) catalysis process. The reason may be that one part of the PS reacted with Fe(0) and oxidized it to Fe(II). Another part of PS reacted with the generated Fe(II) to produce SO₄^•−.

$${\mathrm{S}}_2{\mathrm{O}}_8^{2-}+{\mathrm{SO}}_4^{•-}\to {\mathrm{S}}_2{\mathrm{O}}_8^{•-}+{\mathrm{SO}}_4^{2-}$$

(6)

Figure 4a demonstrates the kinetic models of BPA degradation by Fe(II)/PS at varying PS concentrations. According to the curves in Figure 4a, 1/C_t exhibited a good linear fit to time in the presence of PS with a high linear fit coefficient (R² ≥ 0.87132), indicating that the degradation of BPA followed the pseudo-secondary kinetic model well. The maximum reaction rate of 0.01484 (mg/L)⁻¹min⁻¹ was achieved when the PS concentration was 0.4 mM. The curves in Figure 4b revealed that the degradation of BPA at different PS concentrations in Fe(0)/PS conformed to the pseudo first-order kinetic equation. At a PS concentration of 0.5 mM, the degradation rate of BPA could reach 0.02366 (mg/L)⁻¹min⁻¹. BPA showed different degradation kinetics in Fe(II)/PS vs. Fe(0)/PS systems. This may be caused by the difference in the homogeneous catalysis of Fe(II)/PS and the heterogeneous catalysis of Fe(0)/PS.

2.3. Concentrations of BPA

Figure 5a,b expresses the variation of residual BPA concentration in the solution after degradation of BPA by Fe(II)/PS and Fe(0)/PS at varying initial BPA concentrations, respectively. Figure 5a,b shows the same trend, i.e., the BPA removal effect gradually decreased as the initial BPA concentration increased from 1 mg/L to 5 mg/L. When the initial concentration of BPA was 1 mg/L, the removal rate of BPA was 82.1% in Fe(II)/PS system, and 83.4% in Fe(0)/PS system. When the initial concentration of BPA was increased to 5 mg/L, the removal rate of BPA decreased to 55.3% in Fe(II)/PS system and 45.6% in Fe(0)/PS system. This indicated that the increase in the concentration of the target pollutant (BPA) decreased the BPA degradation efficiency with a constant amount of PS and catalyst (Fe(II) or Fe(0)). Therefore, to obtain higher BPA removal, the amount of catalyst and oxidant needs to be increased proportionally while increasing the initial concentration of BPA.

The kinetic models of BPA degradation were fitted in Figure 6a,b, and the BPA degradation in both systems Fe(0)/PS and Fe(II)/PS conformed to the pseudo first-order kinetic equation (R² ≥ 0.93456). When the initial concentration of BPA was 1 mg/L, the degradation kinetic constant was 0.02537 (mg/L)⁻¹min⁻¹ in Fe(II)/PS system and 0.02854 (mg/L)⁻¹min⁻¹ in Fe(0)/PS system. When BPA was increased to 5 mg/L, the kinetic constant of BPA degradation decreased to 0.01232 (mg/L)⁻¹min⁻¹ in the Fe(II)/PS system, and 0.00994 (mg/L)⁻¹min⁻¹ in the Fe(0)/PS system, which was consistent with the trend of BPA residual concentration in aqueous solution in Figure 5a,b.

2.4. Reaction Temperature

Temperature is an important factor affecting the rate of chemical reactions. The effect of varying temperatures on BPA removal in systems Fe(II)/PS and Fe(0)/PS are given in Figure 7a,b, respectively. With only PS in the system, the degradation rates of BPA in Figure 7a,b all increased with the temperature rising. The degradation of BPA via PS was close to 40% at the temperature of 70 °C in both Figures. According to previous studies [41,42,43], PS could be activated by heat to form SO₄^•− (Equation (7)), which in turn improved the degradation efficiency of organic matter. The elevation in temperature could increase the energy of the reaction system. As this energy exceeded the peroxide bond energy of PS, the peroxide bond broke and SO₄^•− was formed. When both Fe(II) (or Fe(0)) and PS were present in the BPA solution, Fe(II) (or Fe(0)) could activate PS together with the heat energy brought by the high temperature and produce more SO₄^•−. Thus, at 70 °C, the removal of BPA reached 96.7% by Fe(II)/PS and reached 94.5% by Fe(0)/PS.

$${\mathrm{S}}_2{\mathrm{O}}_8^{2-}\stackrel{\hspace{1em}\mathrm{Heat}\hspace{1em}}{\to }2{\mathrm{SO}}_4^{•-}$$

(7)

Figure 8a,b shows the kinetic models of BPA degradation at varying temperature conditions in the Fe(II)/PS system and Fe(0)/PS system, respectively. The pseudo first-order reaction kinetic equations (R² ≥ 0.80025) were applied to the degradation of BPA at different temperatures in both Figure 8a,b. The degradation rates of BPA by the Fe(II)/PS system and Fe(0)/PS system at 70 °C reached 0.06130 (mg/L)⁻¹min⁻¹ and 0.03866 (mg/L)⁻¹min⁻¹, respectively.

In order to assess the relationship between the reaction temperature and degradation rate of BPA, the Arrhenius equation (Equation (8)) [42] was used to fit the experimental data to investigate the correlation.

$$k_{app}=\mathrm{A}\exp (\frac{-{\mathrm{E}}_{\mathrm{a}}}{\mathrm{R}T})\iff \ln k_{app}=\ln \mathrm{A}-\frac{{\mathrm{E}}_{\mathrm{a}}}{\mathrm{R}}\frac{1}{T}$$

(8)

where k_app is the apparent rate constant (min⁻¹), A is the pre-exponential factor (min⁻¹), R is the gas constant (8.314 J/K·mol), T is the absolute temperature (K), and Ea is the activation energy (J/mol). The established thermodynamic equations are shown in Figure 9a,b. In the Fe(II)/PS system, R², A and Ea were 0.94923, 1139.67 min⁻¹, and 2.8 × 10⁷ J/mol, respectively, which showed that the relationship between the BPA degradation rate and the temperature satisfied the Arrhenius equation, while in Fe(0)/PS system, R² was 0.74757. Compared with the homogeneous system Fe(II)/PS, the heterogeneous activation system Fe(0)/PS acquired a low R² when fitting into the Arrhenius equation. This may be because the process of heterogeneous activation involves the release of Fe(0) to Fe(II), which makes the thermodynamic laws difficult to obtain.

2.5. Initial pH

The initial pH value of solution is an important factor affecting the chemical reaction. Therefore, the degradation of BPA by two systems under different initial pH values was studied. Only the initial pH value of the solution was adjusted, and the pH value in the whole reaction process was not controlled. Figure 10a,b reflected the effect of changing pH values on the degradation of BPA by Fe(II)/PS and Fe(0)/PS. It was found that in Figure 10a,b, pH 5.0 was the optimal condition for both systems. The degradation of BPA was poor in both systems under alkaline conditions. On the one hand, the oxidation capacity of the systems was reduced due to the reaction of sulfate radicals converted to hydroxyl radicals in the presence of hydroxide (Equation (9)) [44]. On the other hand, the formation of Fe(OH)₃ could lead to the inhibition of the catalytic reaction (Equation (10)). In the practical application of these two systems, if 5.0 is set as the best pH value, it needs to consume chemical regulators. In addition, the requirements for the corrosion resistance of pipelines have also been improved. One suggestion is that under the condition that partial removal of BPA is allowed in the pretreatment stage, it can be considered to degrade BPA at pH 6.5.

$${\mathrm{SO}}_4^{•-}+{\mathrm{OH}}^{-}\to {\mathrm{SO}}_4^{2-}+{\mathrm{OH}}^{•}$$

(9)

$${\mathrm{Fe}}^{3+}+3{\mathrm{OH}}^{-}\to {\mathrm{Fe}\left(\mathrm{OH}\right)}_3$$

(10)

The kinetic equations for the degradation of BPA by the Fe(II)/PS and Fe(0)/PS systems at different pH values are given in Figure 11a,b, respectively. The BPA degradation curves satisfied the pseudo first-order kinetic equation (R² ≥ 0.89857) except for the degradation of BPA by the Fe(0)/PS system at pH 9. The maximum kinetic constant (0.02818 (mg/L)⁻¹min⁻¹) appeared in the Fe(0)/PS system at pH 5.

2.6. Quenching Experiments and Mechanism Analysis

In order to identify which type of free radicals appeared in the reaction, methanol (Me(OH)) and tert-Butyl alcohol (TBA) were used as extinguishing agents to investigate the degree of inhibition of the chemical reaction. MeOH is an effective scavenger for both •OH (k = 9.7 × 10⁸ M⁻¹ s⁻¹) and SO₄^•− (k = 1.1 × 10⁷ M⁻¹ s⁻¹), while TBA is an effective scavenger for •OH (k = 6.0 × 10⁸ M⁻¹ s⁻¹), but not for SO₄^•− (k = 9.1 × 10⁵ M⁻¹ s⁻¹) [45]. Figure 12a,b exhibits BPA degradation in the presence of either MeOH or TBA in the Fe(II)/PS and Fe(0)/PS systems, respectively. Both MeOH and TBA were found to inhibit BPA degradation in both systems. The inhibitory effect of Me(OH) on the two systems was greater than that of TBA. It suggested that both SO₄^•− and •OH played roles in the oxidation of BPA in the two systems.

Figure 13 depicts the mechanism of PS activation and BPA degradation in the presence of Fe(II) and [Fe(0)] catalysts. In the Fe(0)/PS system, Fe(0) is first transformed into Fe(II) through three pathways (Equations (3)–(5)), thus playing the role of activating PS. In the Fe(II)/PS system, the dissolved Fe(II) directly undergoes a redox reaction with PS to generate SO₄^•− and Fe(III) (Equation (2)). Fe(III) can be transformed into Fe(II) with the gain of an electron. Therefore, Fe(II) and Fe(III) in solution can stably activate PS if they are in a sustainable equilibrium. At the same time, when the OH⁻ concentration in the solution is high, it can react with SO₄^•− to generate hydroxyl radicals, which also has a good oxidation effect on organic matter. Therefore, under the combined action of SO₄^•− and •OH, BPA can be oxidized and the final products will be H₂O and CO₂.

2.7. Comparison of This Study with Other Studies

There are also some previous studies on the degradation of pollutants by Fe(II)/PS and Fe(0)/PS systems. Some examples of representative literature were selected and are listed in

Table 1. Comparison with related process treatment effects in the relevant literature.

System	Reaction Conditions and Pollutant Removal Rate	Ref.
Fe(II)/PS	[acetaminophen (ACT)] = 0.05 mM, [Fe(II)] = 1 mM, [PS] = 0.8 mM, pH = 3.0, T = 20 °C, t = 30 min, 70% ACT removed.	[45]
Fe(II)/PS	[carbamazepine (CBZ)] = 0.025 mM, [Fe(II)] = 0.25 mM, [PS] = 0.25 mM, pH = 3.0, t = 40 min, 57% CBZ removed.	[46]
Fe(II)/PS	[diatrizoate (DTZ)] = 5 mg/L, [Fe(II)] = 1 mM, [PS] = 10 mM, pH = 3.0, T = 25 °C, t = 120 min, 69% DTZ removed.	[47]
Fe(II)/PS	[trimethoprim (TMP)] = 0.05 mM, [Fe(II)] = 4 mM, [PS] = 4 mM, pH = 3.0, T = 25 °C, t = 240 min, 73% TMP removed.	[48]
Fe(II)/PS	[propachlor] = 10 mg/L, [Fe(II)] = 2.5 mM, [PS] = 5 mM, pH = 4.0, T = 30 °C, t = 32 h, 60% propachlor removed.	[49]
Fe(II)/PS	[BPA] = 5 mg/L, [Fe(II)] = 0.1 g/L, [PS] = 0.2 mM, pH = 5.0, T = 20 °C, t = 60 min, 65% BPA removed.	this study
Fe(0)/PS	[p-chloroaniline (PCA)] = 0.05 mM, [Fe(0)] = 0.7 g/L, [PS] = 2.5 mM, pH = 7.0, T = 25 °C, t = 300 min, 65% PCA removed.	[50]
Fe(0)/PS	[aniline] = 0.05 mM, [Fe(0)] = 0.3 g/L, [PS] = 2.5 mM, pH = 7.0, T = 25 °C, t = 300 min, 72% aniline removed.	[51]
Fe(0)/PS	[imidacloprid (IMI)] = 30 mg/L, [Fe(0)] = 1 g/L, [PS] = 5 mM, pH = 7.0, T = 25 °C, t = 60 min, 87% IMI removed.	[52]
Fe(0)/PS	[chloramphenicol (CAP)] = 0.05 mM, [Fe(0)] = 0.125 mM, [PS] = 1 mM, pH = 5.4, T = 25 °C, t = 80 min, 69% CAP removed.	[53]
Fe(0)/PS	[Acid Orange 7 (AO7)] = 30 mg/L, [Fe(0)] = 0.5 g/L, [PS] = 0.3 g/L, pH = 5.8, T = 22 °C, t = 20 min, 38% AO7 removed.	[54]
Fe(0)/PS	[BPA] = 5 mg/L, [Fe(0)] = 0.1 g/L, [PS] = 0.2 mM, pH = 5.0, T = 20 °C, t = 60 min, 85% BPAremoved.	this study

, and they were compared with the research results of this paper. From the perspective of the removal rate of target pollutants, the removal rates obtained in these studies are all concentrated at about 70%. The degradation efficiency of the two systems for BPA in this paper is also comparable to other systems. These target pollutants are recalcitrant organic substances. Thus, to achieve 100% degradation, it is necessary to increase the amount of oxidant and catalyst and control the temperature and pH of the reaction within an appropriate range. During the pretreatment of source water, it can also be considered to control the BPA removal rate at around 70%.

3. Materials and Methods

3.1. Materials

The chemicals used in the experiment are summarized in

Table 2. Chemicals used in this study.

Chemical	Molecular Formula	Provider
Bisphenol A	C15H16O2	McLean Biochemical Technology Co., Ltd., Shanghai, China
Iron powder	Fe	McLean Biochemical Technology Co., Ltd., Shanghai, China
Ferrous sulfate heptahydrate	FeSO4·7H2O	Silong Science Co., Ltd., Shantou, China
Sodium persulfate	Na2S2O8	National Pharmaceutical Group Chemical Reagent Co., Ltd., Shanghai, China
Sodium hydroxide	NaOH	Jindong Tianzheng Fine Chemical Reagent Factory, Tianjin, China
Sulfuric acid	H2SO4	Beijing Chemical Reagents Factory, Beijing, China
Methanol	CH3OH	Chengdu Kolon Chemical Co., Ltd., Chengdu, China
tert-Butyl alcohol	C4H10O	Chengdu Kolon Chemical Co., Ltd., Chengdu, China

. The purity of the chemicals is at least analytically pure. The solutions used in the experiments were prepared using ultrapure water.

3.2. Methods

A BPA stock solution of 1000 mg/L was prepared with ultrapure water and then diluted according to the concentration needed for the experiment. A beaker with a certain concentration of BPA solution was placed on a magnetic stirrer (200 r/min) and stirred for 20 min (to allow the solution to reach the desired initial temperature), and then a certain amount of Fe(II) (or Fe(0)) and PS was added to the beaker to initiate the reaction. At this point, the time was started and samples were taken every 10 min. The sample was filtered through a 0.22 μm filter membrane and injected into a sample vial, to which 0.5 mL of methanol was added to terminate the reaction. The sample vials were put into high performance liquid chromatography (HPLC, G7121A, Agilent Technologies Inc., Beijing, China) to detect the concentration of BPA. Sulfuric acid and sodium hydroxide were used to adjust pH of BPA solution, and different temperatures were achieved by activating the heating button on the magnetic stirrer.

4. Conclusions

In this study, PS activated by the homogeneous catalyst Fe(II) and heterogeneous catalyst Fe(0) were used to degrade the representative endocrine disruptor BPA in water, and the following conclusions were drawn.

• The degradation effect of BPA was similar for both systems (Fe(II)/PS and Fe(0)/PS), and the Fe(0)/PS system had an advantage in the slow release of Fe(II), which was less likely to cause Fe(II) overload resulting in the consumption of free radicals.
• The concentrations of oxidant, catalyst, and substrate had a large influence on the reaction, and the concentration ratio among them was very important for the degradation of BPA.
• Temperature and pH were important factors affecting the reaction. A lower pH (5.0) was very suitable for the degradation of BPA in both systems. High temperature could enhance the degradation of BPA due to the synergistic activation of PS by an iron catalyst and thermal energy.
• Most of the kinetic equations for BPA degradation under different conditions followed pseudo first-order kinetic equations and were highly correlated.