Comparing the suitability of sodium hyposulfite, hydroxylamine hydrochloride and sodium sulfite as the quenching agents for permanganate oxidation

Abstract

An appropriate quenching agent for potassium permanganate (KMnO₄) is necessary for investigating the oxidation kinetics of contaminants by KMnO₄. In this paper, the suitability of three most commonly used inorganic reductants, including sodium hyposulfite (Na₂S₂O₃), hydroxylamine hydrochloride (NH₂OH·HCl) and sodium sulfite (Na₂SO₃), for quenching KMnO₄ was systematically investigated with phenol as a probe contaminant. Na₂S₂O₃ applied with Na₂S₂O₃/KMnO₄ molar ratio of 20.0 was a good choice for quenching KMnO₄ over the pH range of 2.0–11.0. Quenching KMnO₄ with NH₂OH·HCl was a good alternative with 20-fold excess at pH 4.0–9.0. However, Na₂SO₃ was not recommended to be the quenching agent of KMnO₄ due to the significant change of phenol concentration during quenching reaction. The 3-dimensional UV-vis spectra at different pH were collected with stopped flow spectrometer to investigate the quenching mechanisms of these three quenchers toward KMnO₄, which clearly showed the variation of manganese species with time.

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1. Introduction

Potassium permanganate (KMnO₄) has the venerable virtues of easy handling, effectiveness, relatively low cost and comparative stability over a wide pH range.^1–5 Furthermore, no chlorinated or brominated byproducts are formed during its application in the oxidation of natural organic matter.^1,3,4,6,7 Therefore, KMnO₄, as a green oxidant, has been already widely used over the past decades for controlling dissolved Mn(II), taste and odor compounds, cyanotoxins, phenolic compounds, and algae in water treatment plants.^4,7 Although KMnO₄ is widely regarded as a strong oxidant, the oxidation rates of contaminants by KMnO₄ are highly variable depending on the reaction conditions.⁸ Thus, the research of the oxidation kinetics of contaminants by KMnO₄ is critical for comprehending the removal mechanisms of contaminants and predicting the elimination of contaminants in drinking water treatment process under certain conditions.

To determine the oxidation kinetics of contaminants, the samples were collected at selected intervals and quenched before analyzing the concentration of residual contaminants.^9,10 To achieve this objective, reductant is the optimal choice to wipe the oxidants and prevent the continuous decrease of the contaminant concentration. The effects of different quenchers on hydrogen peroxide involved in advanced oxidation process had been investigated, which clarified the necessity of choosing appropriate quenchers to terminate the reaction before the subsequent analysis procedure.^9–11 Liu et al.¹⁰ also reported that the selection of quencher and the amount of quencher were significant to minimize the effects of quenching reaction on the following test and analysis. Different H₂O₂ quenching agents had different influences on subsequent chlorine or chloramine disinfection in uniform formation conditions test and thus the selection of proper quenching agents was crucial.¹⁰

Theoretically, any reagents that can reduce KMnO₄ rapidly can be considered as the possible candidate scavengers for KMnO₄. The ideal reagent will not affect the analysis of residual contaminants after quenching KMnO₄. However, there is no report about the appropriate quenching agent for KMnO₄ over a wide pH range. In this research, three of the most commonly used inorganic reducing reagents, including sodium hyposulfite (Na₂S₂O₃),^2,12 hydroxylamine hydrochloride (NH₂OH·HCl)¹³ and sodium sulfite (Na₂SO₃),¹⁴ which had been used as quenchers by researchers when they investigated the kinetics of contaminants oxidation by KMnO₄, were selected to evaluate their suitability as quenching agent for KMnO₄. Our previous studies demonstrated that Na₂SO₃ enhanced the oxidation rate and oxidizing ability of KMnO₄ significantly with the Na₂SO₃/KMnO₄ molar ratio of 5 : 1 at pH 4.0–9.0 ¹⁵ and Na₂S₂O₃ accelerated the oxidation of phenol by KMnO₄ with Na₂S₂O₃/KMnO₄ molar ratio of 1 : 5 at pH 5.0.¹ However, the suitability of Na₂SO₃ and Na₂S₂O₃ as quenchers for KMnO₄ at high Na₂SO₃/KMnO₄ and Na₂S₂O₃/KMnO₄ molar ratios keeps unknown.

Many phenols such as phenol, 2-chlorophenol (2-CP), 2,4-dichlorophenol (2,4-DCP) and trichlorophenols have been designated as the priority pollutants by the U.S. Environmental Protection Agency (EPA) since 1979.¹ Some environmentally topical endocrine disrupting chemicals (EDCs) which contain phenolic hydroxyl group such as bisphenol A (BPA), triclosan (TCS), estrone (E1), 17β-estradiol (E2), estriol (E3), 2,4-DCP and 4-n-nonylphenol (4-n-NP), have been frequently detected in surface waters.^16,17 Therefore, phenol was selected as a probe contaminant in this study. The objective of this study was to (1) determine the appropriate quenching agent and quenching agent dose for KMnO₄ over a wide pH range; (2) investigate the reaction mechanisms of different quenchers with KMnO₄.

2. Experimental section

2.1. Materials

KMnO₄ (guaranteed reagent (GR) grade), Na₂S₂O₃·5H₂O (GR grade), and phenol of 99% purity were purchased from Tianjin Chemicals Reagent Co., Ltd. (Tianjin, China). Aniline (99% pure) was supplied by Sigma-Aldrich (St. Louis, MO, USA). Na₂SO₃ with the purity of 97% and NH₂OH·HCl of 98.5% purity were obtained from Qiangsheng Jiangsu and Guoyao China, respectively. Methanol (HPLC grade), acetonitrile (HPLC grade) and formic acid (GR grade) were purchased from Merck (Darmstadt, Germany) for the chemical analysis. All chemicals were used without further purification and all solutions were prepared in 18.2 MΩ cm Milli-Q water from a Millipore system.

The KMnO₄ crystal was dissolved in Milli-Q water to prepare a 50 mM stock solution. The stock solution of Na₂S₂O₃ (250 mM), NH₂OH·HCl (250 mM) and Na₂SO₃ (250 mM) were freshly prepared for each set of experiments to avoid oxidation by oxygen. The stock solutions of phenol (1.0 mM) and aniline (1.0 mM) were prepared in Milli-Q water every day.

2.2. Batch experiments

In this study, 100 mL volumetric flasks (batch reactors) were used to perform the batch oxidation experiments at room temperature (25 ± 1.0 °C). To determine the kinetics of phenol oxidation at various pH levels, the reaction was initiated by quickly spiking excessive KMnO₄ into the solutions containing phenol at different pH. Aliquots were then periodically withdrawn and quenched with more than 2 mM of Na₂S₂O₃. Quenching reactions were initiated by quickly adding the stock KMnO₄ solution into the reactor containing phenol and quencher of various concentrations and the reaction time was 5 h. Then the sample was filtered with 0.22 μm membrane, and collected into sample vials for subsequent phenol analysis with ultra-performance liquid chromatography (UPLC). All experiments were performed at least in triplicate.

2.3. Stopped-flow experiments

A stopped-flow spectrophotometer (SFS, Model SX20, Applied Photophysics Ltd., Leatherhead, UK) was used to investigate the reduction process of KMnO₄ by different quenchers. A photodiode array for acquisition of multi-wavelength absorption and a fluorimeter were used as the detectors with a 150 W xenon lamp as the light source. An HP computer workstation was employed to control the stopped-flow and acquire the kinetic data.

For the stopped-flow experiments, an equal volume of KMnO₄ and reductants (in the presence/absence of phenol) solution were simultaneously injected into the optical cell of the SFS with two automatic syringes driven by compressed nitrogen. The solutions were adjusted to the target pH levels by adding HCl or NaOH before injecting into the SFS. The 3-dimensional UV-visible spectrums of the reaction between quenchers and KMnO₄ were conducted using photodiode array at 350–700 nm and the variation of phenol concentration was continuously detected by fluorimetry at E_x/E_m = 272 nm/298 nm.¹⁸ All stopped-flow experiments were performed at least in duplicate and the average value was reported.

2.4. Chemical analysis

A high performance Leici pH meter with a saturated potassium chloride solution as electrolyte (Shanghai, China) was used to measure solution pH. Daily calibration with proper buffer solutions (pH 4.00, 6.86 and 9.18) was performed to ensure its accuracy. A UPLC (Waters ACQUITY UPLC H-Class) consisted of a quaternary solvent manager (QSM), a sample manager (FTN), fluorescence detector (FLR) and an UV detector (TUV) was used for analyzing the concentration of aniline and phenol (Milford MA, United States). Separation was accomplished with an UPLC BEH C18 column (2.1 × 50 mm, 1.7 μm; Waters) at 35 ± 1.0 °C and a mobile phase of methanol-0.1% formic acid aqueous solution (60 : 40) for phenol and acetonitrile–water (30 : 70) for aniline. The flow rate was 0.3 mL min⁻¹ and the largest injection volume was 10 μL. The concentration of phenol was determined with TUV by comparing the peak area at 273 nm with the standard solution of phenol and that of aniline was determined with FLR by comparing the peak area at E_x/E_m = 232 nm/329 nm with the standard solution of aniline.

3. Results and discussion

3.1. Influence of pH on oxidation kinetics of phenol by KMnO₄

The mechanisms of organics oxidation by KMnO₄ are different in different pH regions, i.e., acid-catalyzed reaction below pH 5.0, uncatalyzed reaction in the pH range of 6.0–9.0, and base-catalyzed reaction above pH 10.0.¹⁹ To clarify the performance of different reductants for quenching KMnO₄ at different pH, the time courses of oxidative degradation of phenol by KMnO₄ in 10-fold excess over the pH range of 2.0–11.0 were investigated and shown in Fig. 1. Auto-acceleration was observed in the process of phenol oxidation by KMnO₄ at acidic pH. As shown in Fig. S1,† lag phase was also observed during aniline oxidation by KMnO₄ at pH 4.0, suggesting the ubiquity of acid-catalysis phenomenon in KMnO₄ oxidation at acidic pH. Previous studies showed that acid-catalysis of KMnO₄ oxidation might be ascribed to the generation of Mn(II),²⁰ the in situ formed manganese dioxide (MnO₂)²¹ and other manganese intermediates.²² We fit the majority of the data with pseudo-first-order kinetics after excluding a few data points at the beginning of the reaction and summarized the oxidation rate constants of phenol by KMnO₄ in Table S1.† The minimum oxidation rate of phenol by KMnO₄ was obtained at pH 6.0 and the rate constants of phenol oxidation by KMnO₄ showed a monotonic increase with increasing pH from 6.0 to 10.0, and then dropped with further increase in pH when the dissociated phenol predominated in solution. The variation of oxidation rate of phenol with pH was consistent with our previous study which showed that the undissociated phenol could be directly oxidized by KMnO₄ while H⁺ was combined during the oxidation of dissociated phenol.²

Fig. 1 Time course of phenol oxidation by permanganate at pH values ranging from 2.0 to 11.0. Reaction conditions: [phenol]₀ = 5 μM, [KMnO₄]₀ = 50 μM.

3.2. Quenching with Na₂S₂O₃

A good quenching agent could exhaust KMnO₄ immediately and minimize the elimination of phenol by KMnO₄ over a wide pH range. Therefore, the ratio of residual concentration of phenol to the initial phenol concentration (C/C₀) was employed to evaluate the effectiveness of the KMnO₄ quenchers. The ideal value of C/C₀ was 1.0 when the quenching agent was very efficient. Firstly, the performance of Na₂S₂O₃ of different concentrations in quenching KMnO₄ at pH 2.0–11.0 was examined and shown in Fig. 2. Preliminary experiments in this study had excluded the interaction between Na₂S₂O₃ and phenol at pH 2.0–11.0. Thus, the drop in phenol concentration should be ascribed to its oxidation by KMnO₄ or the in situ formed manganese intermediates.

Fig. 2 Quenching of residual permanganate with Na₂S₂O₃ of different concentrations in phenol oxidation. Reaction conditions: [phenol]₀ = 5 μM, [KMnO₄]₀ = 50 μM.

Depending on pH, Na₂S₂O₃ can be oxidized by KMnO₄ to sulfate through one of the following reaction routes, as illustrated in eqn (1) and (2).

5S₂O₃²⁻ + 8MnO₄⁻ + 7H⁺ → 8Mn²⁺ + 10SO₄²⁻ + 7OH⁻ (1)

3S₂O₃²⁻ + 8MnO₄⁻ + H₂O → 8MnO₂ + 6SO₄²⁻ + 2OH⁻ (2)

Eqn (1) and (2) revealed that 50 μM Na₂S₂O₃ was enough to reduce 50 μM KMnO₄ to either Mn²⁺ or MnO₂. However, when Na₂S₂O₃ was applied at 50 μM, the concentration of phenol remained in the solution was always lower than its initial concentration at pH 2.0–11.0, indicating that the reduction of KMnO₄ to stable manganese species by Na₂S₂O₃ was not fast enough to avoid the subsequent oxidation of phenol. As the concentration of Na₂S₂O₃ was increased from 50 μM to 200 μM, the quenching efficiency of Na₂S₂O₃ increased significantly over the pH range of 2.0–11.0. However, the variation of quenching efficiency with pH at 200 μM Na₂S₂O₃ was similar to that at lower Na₂S₂O₃ concentration. The quenching efficiencies of Na₂S₂O₃ were as high as 93.0–100%, 95.1–100%, and 96.8–100%, respectively, over the pH range of 2.0–11.0 when the concentration of Na₂S₂O₃ was increased to 500, 1000 and 2000 μM. Therefore, Na₂S₂O₃ dosed at ≥1000 μM was recommended to quench 50 μM KMnO₄ to ensure the precise determination of residual phenol over the pH range of 2.0–11.0.

Fig. 3 shows the online scanning of 3D UV-vis spectra at 350–700 nm (where phenol and its oxidation products do not absorb) during the reactions between KMnO₄ and Na₂S₂O₃ with the molar ratio of 1 : 10 over the pH range of 2.0–11.0. KMnO₄ was reduced to Mn(II) directly at pH 2.0–3.0 while colloidal MnO₂ with the nonspecific absorbance at <500 nm ¹ was generated rapidly at pH 4.0–11.0. Hereafter, the generated MnO₂ at pH 4.0–10.0 was further reduced gradually to Mn(II) with higher reaction rate at lower pH. At pH ≥ 10.0, the generated colloidal MnO₂ was stable even in the presence of excess Na₂S₂O₃. Comparing the generation and consumption of MnO₂, the reaction rate between Na₂S₂O₃ and KMnO₄ was much greater than that between Na₂S₂O₃ and MnO₂ at pH 4.0–11.0. In fact, MnO₂ might be generated and reduced with higher reaction rate at pH 2.0–3.0 as MnO₂ is a strong oxidizing agent under acidic conditions, which accounted for the failure in detecting the in situ MnO₂.

Fig. 3 The evolution of three dimensional UV-visible spectra during the reaction between Na₂S₂O₃ and KMnO₄ over the pH range of 2.0–11.0. Reaction conditions: [KMnO₄]₀ = 50 μM, [Na₂S₂O₃]₀ = 500 μM, [phenol]₀ = 5 μM.

To further clarify the quenching mechanism of Na₂S₂O₃ toward KMnO₄, the rate constants of KMnO₄ reduction by Na₂S₂O₃ was evaluated. The loss of KMnO₄, characterized by the drop in the absorbance at 525 nm shown in Fig. 3, followed pseudo-first-order kinetics with Na₂S₂O₃ in 10-fold excess at pH 2.0–11.0, suggesting that the reduction of KMnO₄ by Na₂S₂O₃ was first-order with respect to KMnO₄. The obtained rate constants of KMnO₄ reduction by Na₂S₂O₃, summarized in Table S1,† dropped with increasing pH from 2.0 to 5.0 and almost keep constant at pH ≥ 6.0. The higher reduction rate of KMnO₄ by Na₂S₂O₃ at acidic pH could be attributed to the acid-catalyzed reaction and the higher oxidation reduction potential of KMnO₄. In spite of the comparatively low reaction rate of Na₂S₂O₃ with KMnO₄ at pH ≥ 6.0, the reduction rate of KMnO₄ by Na₂S₂O₃ was still much higher than that by phenol at the same pH level (as shown in Table S1†), which accounted for the satisfactory quenching efficiency.

Our previous study reported that Na₂S₂O₃ dosed in the process of phenol oxidation by KMnO₄ at Na₂S₂O₃/KMnO₄ molar ratio of 1 : 5 would induce the formation of colloidal MnO₂, which oxidized phenol at a higher rate than KMnO₄ at pH 5.0 and thus the application of Na₂S₂O₃ resulted in a faster phenol degradation by KMnO₄.¹ However, efficient quenching was achieved at pH 5.0 by adding 10-fold excess Na₂S₂O₃, which should be attributed to the much higher reduction rate of MnO₂ by Na₂S₂O₃ than by phenol. As pH highly influenced the activity of MnO₂, satisfactory quenching efficiency was achieved at high pH even when intermediate MnO₂ co-existed. The above results suggested that Na₂S₂O₃ was a proper quencher for the rapid reactions between contaminants and KMnO₄. But, Na₂S₂O₃ was not proper for quenching the reaction involving contaminants which hold strong interactions with MnO₂.

3.3. Quenching with NH₂OH·HCl

The performance of NH₂OH·HCl as KMnO₄ quenching agent over the pH range of 2.0–11.0 was also examined and shown in Fig. 4. It was found that NH₂OH·HCl dosed at 50 μM had very poor quenching performance toward KMnO₄ since the concentration of residual phenol was only 0–84.1% of its initial concentration. As can be seen, the variation of the quenching efficiency of 50 μM NH₂OH·HCl with pH was very different from that of Na₂S₂O₃ applied at the same concentration, indicating that these two quenching agents reduced KMnO₄ with different mechanisms. Increasing NH₂OH·HCl concentration from 50 μM to 200 μM, the quenching efficiency at pH 2.0–11.0 elevated considerably and could reach 91.8–97.9% at pH 4.0–9.0. When the concentration of NH₂OH·HCl was further increased to 500 μM, 1000 μM, and 2000 μM, the quenching efficiencies of NH₂OH·HCl were 92.0–99.7%, 95.0–98.7%, 97.5–98.5% at pH 4.0–9.0, respectively. However, at pH level as high as 10.0–11.0, the quenching efficiency of NH₂OH·HCl was always lower than 85.5% even when it was dosed in large excess (500–2000 μM). At pH 2.0–3.0, the quenching efficiency of NH₂OH·HCl enhanced progressively with increasing NH₂OH·HCl concentration. But the quenching efficiency was ≤92.5% except the case when NH₂OH·HCl was dosed at 2000 μM at pH 3.0. In sum, NH₂OH·HCl applied with NH₂OH·HCl/KMnO₄ molar ratio of 20 can work as KMnO₄ quenching agent at pH 4.0–9.0 to achieve a quenching efficiency of ≥95.0%. However, NH₂OH·HCl was not recommended to quench KMnO₄ at pH < 4.0 or pH > 9.0.

Fig. 4 Quenching of residual permanganate with NH₂OH·HCl of different concentrations in phenol oxidation. Reaction conditions: [phenol]₀ = 5 μM, [KMnO₄]₀ = 50 μM.

Fig. 5 showed the 3D UV-vis spectra at 350–700 nm collected during the reaction between KMnO₄ and NH₂OH·HCl with the KMnO₄/NH₂OH·HCl molar ratio of 1 : 10 over the pH range of 2.0–11.0. At pH 2.0–6.0, no MnO₂ was observed in the process of KMnO₄ reduction by NH₂OH·HCl. At pH 7.0–11.0, MnO₂ was formed in situ and then gradually reduced. The accumulation of MnO₂ might be due to the decrease of MnO₂ activity and the increase of NH₂OH·HCl redox potential with increasing pH (Fig. S2†). The inefficiency of NH₂OH·HCl dosed at NH₂OH·HCl/KMnO₄ molar ratio of 20 for quenching KMnO₄ at pH 2.0–3.0 should be mainly associated with the low reduction rate of KMnO₄ by NH₂OH·HCl, as shown in Table S1.† The reaction rate constant of KMnO₄ with NH₂OH·HCl was lower than that of KMnO₄ with phenol at pH 2.0 and was only 7.4-flod of that of KMnO₄ with phenol at pH 3.0. The reduction rate of KMnO₄ decreased with increasing pH from 3.0 to 4.0, increased significantly with increasing pH from 5.0 to 8.0 and kept stable at pH 9.0–11.0 (Table S1†). Considering that the pK_a of NH₂OH is 5.9,²³ the specie of NH₃OH⁺ gradually transformed to NH₂OH as pH increased from 5.0 to 7.0. The oxidation reduction potential of NH₂OH was much lower than NH₃OH⁺ (as shown in Fig. S2†), which accounted for the increased rate of KMnO₄ reduction with increasing pH from 5.0 to 8.0. Strangely, NH₂OH·HCl was not effective to be a quenching agent for KMnO₄ at pH 10.0 and 11.0 in spite of the high reaction rate between KMnO₄ and NH₂OH·HCl, which need further investigation.

Fig. 5 The evolution of three dimensional UV-visible spectra during the reaction between NH₂OH·HCl and KMnO₄ over the pH range of 2.0–11.0. Reaction conditions: [KMnO₄]₀ = 50 μM, [NH₂OH·HCl]₀ = 500 μM, [phenol]₀ = 5 μM.

3.4. Quenching with Na₂SO₃

Although Na₂S₂O₃ accelerated the oxidation of phenol by KMnO₄ with Na₂S₂O₃/KMnO₄ molar ratio of 1 : 5 at pH 5.0,¹ this study showed that Na₂S₂O₃ dosed at Na₂S₂O₃/KMnO₄ molar ratio of 20 : 1 was efficient to quench KMnO₄ over the pH range of 2.0–11.0. Therefore, it is necessary to investigate the quenching efficiency of Na₂SO₃ with different concentrations over a wide pH range although our previous study reported the acceleration of phenol oxidation by KMnO₄ in the presence of Na₂SO₃ at Na₂SO₃/KMnO₄ molar ratio of 5 at pH 4.0–9.0.¹⁵

Na₂SO₃ had also been employed as KMnO₄ quenching agent in the literature.¹⁴ However, the results of this study showed that it was a very poor quencher for KMnO₄ at Na₂SO₃/KMnO₄ molar ratio ranging from 1 to 40, especially at pH < 7.0 (Fig. 6). The concentration of phenol decreased rapidly during its oxidation by KMnO₄ in the presence of Na₂SO₃ at Na₂SO₃/KMnO₄ molar ratio of 10 at pH 5.0, as shown in Fig. S3,† which should be ascribed to the generation of highly active Mn(III), as reported in our previous study.¹⁵

Fig. 6 Quenching of residual permanganate with Na₂SO₃ of different concentrations in phenol oxidation. Reaction conditions: [phenol]₀ = 5 μM, [KMnO₄]₀ = 50 μM.

Similar to Na₂S₂O₃ and NH₂OH·HCl, the quenching performance of Na₂SO₃ became better with increasing concentration over the pH range of 2.0–11.0 (Fig. 6). However, Na₂SO₃ had much worse quenching performance than Na₂S₂O₃ and NH₂OH·HCl. The largest quenching efficiency of Na₂SO₃ applied at 50–2000 μM at pH 2.0–11.0 was 95.8%, which was achieved at pH 9.0 with 2000 μM Na₂SO₃. However, the quenching efficiency was as low as 0–92.4% under other conditions investigated in this study.

To characterize the overall reactions involved in the KMnO₄/Na₂SO₃ process, the variation of UV-vis absorption at 350–700 nm during the reactions between KMnO₄ and Na₂SO₃ with the molar ratio of 1 : 10 over the pH range of 4.0–11.0 was shown in Fig. 7. At pH 2.0 and 3.0, the rate of KMnO₄ reduction by Na₂SO₃ to Mn(II) was too fast to be detected by SFS. At pH 6.0, the spectrum of KMnO₄ was dominant at the beginning, but its strong absorbance at 300–350 and 500–570 nm disappeared at about 50 ms. At 200 ms, an obviously broad absorbance shoulder developed at <500 nm (characteristic of colloidal MnO₂).⁵ This trend was similar to that observed in our previous study,²⁴ suggesting a rapid but multi-step reduction of KMnO₄ to MnO₂ via Mn(III) intermediates. Mn(III) is labile and susceptible to disproportionate to Mn(II) and Mn(IV) due to the tetragonal distortion of electron configuration of Mn(III).²⁵

Fig. 7 The evolution of three dimensional UV-visible spectra during the reaction between Na₂SO₃ and KMnO₄ over the pH range of 2.0–11.0. Reaction conditions: [KMnO₄]₀ = 50 μM, [Na₂SO₃]₀ = 500 μM, [phenol]₀ = 5 μM.

The rapid disappearance of phenol suggested the higher reaction rate of Mn(III) towards phenol than its disproportionation and reduction by bisulfite/sulfite. The second-order rate constants of contaminants' oxidation by Mn(III) could be obtained by constructing the kinetic model, as illustrated in our previous study.²⁴ To calculated the second-order rate constants of phenol oxidation by Mn(III), relative rate method was employed. Bisphenol A (BPA) was selected as the reference compound since the second-order rate constants of its reaction with Mn(III) had been obtained.²⁴ The generated Mn(III) reacts with BPA and phenol, with second-order rate constants k_BPA and k_phenol, respectively:

The degradation kinetics of BPA and phenol can be expressed as follows:

Rearranging and integrating eqn (5) and (6), we obtains:

Thus,

Therefore, a plot of vs. should be a straight line with the slope of , as shown in Fig. S4.† By changing the initial concentration of BPA and phenol, the rate constants k_phenol were calculated and summarized in Table S2.† The rate constants of the reaction between phenol and Mn(III) was as high as 10⁴ to 10⁵ M⁻¹ s⁻¹ at pH 5.0 and 7.0, thus the quenching efficiency of Na₂SO₃ was low. At pH ≥ 7.0, more MnO₂ was generated, as shown in Fig. 7, due to the higher disproportionation rate of Mn(III) than the reduction rate of Mn(III) by sulfite.²⁶ The higher disproportionation rate of Mn(III) decreased the utilization of Mn(III) by phenol under alkaline condition, and thus enhanced the quenching efficiency of KMnO₄ by sulfite.

4. Conclusions

In this paper, the quenching efficiency of KMnO₄ by Na₂S₂O₃, NH₂OH·HCl and Na₂SO₃ were systematically investigated over a wide pH range and a wide reductant/KMnO₄ molar ratios. The results show that Na₂S₂O₃ is the viable option for various pH values. NH₂OH·HCl is another good choice for quenching KMnO₄ with 20-fold excess at pH 4.0–9.0, while not recommended at pH < 4.0 or >9.0. pH adjustment of quenching solution is a feasible method to expand the pH range of quenching application. The use of Na₂SO₃ is not recommended for quenching KMnO₄ over a wide pH range due to the significant change of phenol concentration during quenching reaction.

The reduction mechanisms of KMnO₄ by quenchers were different, depending on the species of quenchers and pH. For Na₂S₂O₃, KMnO₄ was reduced to Mn(II) at pH 2.0–3.0 with 10-fold excess of Na₂S₂O₃. Over the pH range of 4.0–9.0, KMnO₄ was reduced to MnO₂ followed by the further reduction to Mn(II). While at pH 10.0–11.0, KMnO₄ was reduced to stable MnO₂ by Na₂S₂O₃. KMnO₄ was reduced to Mn(II) with 10-fold excess of NH₂OH·HCl over the pH rang of 2.0–6.0. At pH 7.0–11.0, MnO₂ was formed rapidly and gradually reduced to Mn(II). For Na₂SO₃, KMnO₄ was reduced Mn(III), and then Mn(III) was reduced by phenol or excess Na₂SO₃ to Mn(II) or disproportionated to Mn(II) and MnO₂.

Acknowledgements

This work was supported by the Major Science and Technology Program for Water Pollution Control and Treatment (Grant 2012ZX07403-001), the National Natural Science Foundation of China (Grant 21522704) and the Fundamental Research Funds for the Central Universities. We thank Shanghai Institute of Organic Chemistry (Chinese Academy of Sciences) for offering access to the SFS.

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Footnote

† Electronic supplementary information (ESI) available. See DOI: 10.1039/c6ra01209d

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