Auto-accelerating and auto-inhibiting phenomena in the oxidation process of organic contaminants by permanganate and manganese dioxide under acidic conditions: effects of manganese intermediates/products

Abstract

Considering the confused/controversial kinetics law of organic contaminant oxidation by permanganate (MnO₄⁻) and manganese dioxide (MnO₂) under acidic conditions, this study was conducted to systematically investigate the oxidation kinetics and mechanisms of contaminants by MnO₄⁻ and MnO₂ under acidic conditions. The process of phenol oxidation by MnO₄⁻ showed an auto-accelerating trend at pH 2.0–6.0, which was mainly associated with the generation of MnO₂, an oxidant more active than MnO₄⁻ under acidic conditions. However, an auto-inhibiting trend was observed during phenol oxidation by MnO₂, which could be ascribed to the generation of Mn(II) and its subsequent adsorption on the surface of MnO₂. The presence of excessive pyrophosphate (PP) greatly accelerated the oxidation of phenol by MnO₄⁻ and MnO₂ under acidic conditions due to the generation and accumulation of highly reactive Mn(III)–PP. The presence of PP also changed the manganese species at equilibrium from MnO₂ and Mn(II) to Mn(III)–PP, respectively, in the processes of organic contaminant oxidation by MnO₄⁻ and MnO₂. In short, the reduction products of MnO₄⁻ and MnO₂ determined the auto-accelerating or auto-inhibiting phenomenon observed in the process of phenol oxidation by MnO₄⁻ and MnO₂.

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1. Introduction

Water is valuable and crucial to all living organisms and for multiple human activities.¹ However, various synthetic organic compounds, used by society in vast quantities for a range of purposes, have been detected in water with negative impact on water quality.^1–8 Oxidation is a widely used method for removing organic contaminants from water. Various chemical oxidants, such as chlorine,⁹ ozone,¹⁰ ferrate,¹¹ chlorine dioxide,¹² Fenton reagent¹³ and permanganate (MnO₄⁻),¹⁴ can be applied for the transformation/elimination of the undesired organic contaminants. Chlorine is one of the most commonly used disinfectants for drinking water and wastewater treatment.¹⁵ Unfortunately, chlorination generates chlorinated disinfection byproducts which are associated with cancer, particularly bladder and rectal cancer.^16,17 The application of ozone in water treatment is also widespread for disinfection and oxidation.¹⁸ However, ozonation suffers from the potential formation of the carcinogenic brominated byproducts when bromide is present in water.^19,20 Ferrate is an environmentally friendly oxidant which can effectively remove the organic contaminants containing electron-rich moieties.^21,22 Nonetheless, ferrate is very instable in water and difficult to be prepared and stored, which limits its field treatment application. The application of chlorine dioxide has the risk of chlorate generation despite of its efficient transformation of various contaminants.²³ By generating highly reactive hydroxyl radicals, Fenton reagent can eliminate a variety of organic pollutants with different moeities. However, the application of Fenton reagent is limited by its narrow applicable pH range (2.0–4.0) and the negative effects of HCO₃⁻ and humic acid, which are commonly present in water and can consume hydroxyl radical and decrease the effectiveness of Fenton oxidant.²⁴ Another alternative technology for the oxidative removal of organic contaminants is the application of manganese oxidants, which can address many concerns of above-mentioned methods.²⁵

Compared with other oxidants, MnO₄⁻, as a green oxidant,²⁶ is preferred for its attractive characteristics of easy handling, relatively low cost, effectiveness, comparative stability over a wide pH range, and non-formation of halogenated byproducts.²⁷ It has been already widely used for controlling dissolved manganese, taste/odor/color, and biological growth.^25,28–30 MnO₄⁻ has also been proved to be a promising technology for oxidative removal of various phenolic compounds (bisphenol A, triclosan, estrone, 17β-estradiol, estriol, chlorophenols, bromophenols, etc.),^31–35 and other organic chemicals that contain electron-rich moieties, including olefin, amino, thiol, ether, aldehyde, and ketone groups.^25,36 Another common manganese oxidant, manganese dioxide (MnO₂) has been demonstrated to be efficient in degrading various organic compounds including antibiotics,^37,38 anilines,³⁹ phenol,⁴⁰ steroid estrogens,⁴¹ etc. In addition, MnO₂ can act as coagulant aid to improve the performance of coagulation and filtration process or as adsorbent to further remove the constituents of concern.^29,40,42

Although manganese oxides have been widely applied as oxidants, the oxidation kinetics of organic contaminants by manganese oxides are variable under different conditions.^34,35 Generally, the pH-rate profile of MnO₄⁻ oxidation has three regions including an acid-catalyzed reaction at pH ≤ 5.0, an uncatalyzed reaction in the pH range of 6.0–9.0, and a base-catalyzed reaction at pH ≥ 10.0.⁴³ However, the reaction mechanisms of MnO₄⁻ oxidation under acidic conditions remain controversy or obscure. Bahrami et al. showed the delayed auto-catalytic kinetics of glycine, L-alanine and L-leucine oxidation by MnO₄⁻ in strong acidic media, which was ascribed to the accumulation of Mn(II).⁴⁴ Perez-Benito et al. also reported the auto-catalytic oxidation of glycine, L-alanine and L-leucine by MnO₄⁻ but they concluded that the reactions were auto-catalyzed by colloidal MnO₂.⁴⁵ Jáky et al.³⁷ and Wiberg et al.⁴⁶ observed that oxidation of acetylacetone and aromatic aldehydes by MnO₄⁻ was acid-catalyzed in acidic media. Pimienta et al.⁴⁷ elaborated that the reduction of MnO₄⁻ by oxalic acid in sulfuric acid medium could be well described by a model incorporating the specific reactivities of MnO₄⁻ and of various Mn(III) and Mn(IV) reaction intermediates. Compared to MnO₄⁻, less information on the kinetics of organic contaminants oxidation by MnO₂ was available. It is well known that pH significantly influences the activity of MnO₂ towards contaminants with higher reaction rate at lower pH.^38,48 However, the rate law of contaminants degradation by MnO₂ was unclear under different conditions which deserved further study. In water and wastewater, many constituents, such as pyrophosphate (PP), phosphate, ethylene diamine tetraacetic acid (EDTA), and humic acid, can act as ligands to stablize Mn(III) and thus may change the balance of manganese species and influence the oxidation of contaminants accordingly.³² Therefore, it is also necessary to investigate the degradation law of contaminants by manganese oxides in the presence of ligands.

In short, the objectives of this study were to investigate the kinetics of organic contaminant oxidation by MnO₄⁻ and MnO₂ under acidic conditions and explore the mechanisms behind the different kinetic behaviors. Phenol was selected as the major probe contaminant in this study because of the wide existence of phenolic hydroxyl group in some environmentally topical endocrine disrupting chemicals that have been frequently detected in surface waters.^49,50

2. Experimental section

2.1. Materials

Potassium permanganate (GR grade), sodium thiosulfate pentahydrate (GR grade) and phenol (99% pure) were purchased from the Tianjin Chemical Reagent Co., Ltd. (Tianjin, China). Sodium pyrophosphate (PP, AR grade) was supplied by the Sinopharm Chemical Reagent Co., Ltd. (Shanghai, China). Aniline (99% pure) was purchased from Sigma-Aldrich (St. Louis, MO, USA), and methanol (99.9% pure) was supplied by Merck KgaA (Germany). All chemicals were used as received.

The KMnO₄ crystal was dissolved in Milli-Q water to make a 50 mM stock solution. The stock solutions of phenol (1.0 mM), Na₂S₂O₃ (100.0 mM), and Na₄P₂O₇ (50.0 mM) were prepared in Milli-Q water every day. A stable colloidal MnO₂ stock solution was prepared freshly before use following the procedure in the literature⁵¹ by mixing the appropriate amounts of MnO₄⁻ and Na₂S₂O₃ stock solutions.

The stock solution of Mn(III)–PP was prepared following the procedure reported in the literature.^52,53 In brief, sodium pyrophosphate decahydrate was dissolved in water to reach a 50 mM PP final concentration and the pH was adjusted to 8.2 with HCl/NaOH, then manganese(III) acetate dihydrate solid was slowly added to the PP solution during vigorous stirring to reach a final concentration of 10 mM Mn(III). pH of the obtained solution was adjusted to 7.0 to ensure the stability of working solution against disproportionation.

2.2. Batch experiments

The kinetic experiments were conducted in 0.25 L brown glass bottles at 25.0 ± 1.0 °C. Sodium acetate (1 mM) was used as a buffer for the reactions at pH ≤ 6.0, while sodium borate (10 mM) was employed for those at pH ≥ 8.0. However, no buffer was used for the experiments conducted at pH 7.0, which was kept constant at 7.0 ± 0.1 during the reactions with the addition of HCl or NaOH if necessary. The negligible effects of acetate and borate buffers on these reactions have been discussed previously.^34,54,55 Reactions were initiated by quickly spiking excess KMnO₄, MnO₂ or Mn(III)–PP into the solutions containing phenol/aniline while they were being stirred. The total reaction time was determined depending on the removal of contaminant. Periodically, 15 mL of sample was rapidly transferred into a 25 mL beaker, immediately quenched with 100 μL of a sodium thiosulfate stock solution,⁵⁶ filtered with a 0.22 μm membrane, and quickly collected into sample vials for subsequent analysis. All experiments were performed at least in duplicate and the average value was reported unless otherwise noted.

2.3. Chemical analysis

A pH meter with a saturated KCl solution as the electrolyte 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. The concentrations of phenols and aniline were analyzed by an ultraperformance liquid chromatograph (waters ACQUITY UPLC H-Class), consisted of quaternary solvent manager (QSM), a sample manager (FTM) and a UV-visible detector (TUV). Separation was accomplished with a UPLC BEH C 18 column (2.1 × 50 mm, 1.7 μm; Waters) at 35 ± 1 °C. The flow rate was 0.4 mL min⁻¹ and the largest volume injection was 10 μL. The mobile phase of methanol–0.1% formic acid aqueous solution and methanol aqueous solution were used to for phenol and aniline separation, respectively. Concentrations of phenol and aniline were determined by comparing the peak area at 273 nm and Ex/Em = 232 nm/329 nm with that of the corresponding standards, respectively. The spectral changes during the oxidation of phenol by MnO₄⁻ were monitored in the range of 250–800 nm in a thermostated cell compartment using Purkinje TU-1902 automatic scanning UV-vis spectrophotometers with wavelength program controllers.

3. Results and discussion

3.1. The auto-accelerating phenomenon in contaminants oxidation by MnO₄⁻ under acidic conditions

Fig. 1 showed the time course of phenol oxidation by MnO₄⁻ over the pH range of 2.0–9.0, where the initial concentrations of phenol and MnO₄⁻ were 5 and 50 μM, respectively. Obviously, the oxidation kinetics of phenol displays auto-acceleration at pH ≤ 6.0 with higher oxidation rate at lower pH. To verify that the auto-accelerating behavior of MnO₄⁻ oxidation under acidic conditions was not limited to phenol, the kinetics of aniline oxidation by MnO₄⁻ at pH 4.5 was also determined, as demonstrated in Fig. S1.† The oxidation of aniline by MnO₄⁻ at pH 4.5 also exhibited the characteristics of auto-acceleration. Furthermore, the auto-accelerating behavior had been reported in the oxidation of triclosan,³¹ Mn(II),⁵⁷ α-amino acids, acetylacetone,⁵⁸ and aromatic aldehydes⁵⁹ by MnO₄⁻. Thus, the auto-accelerating reaction was irrelevant to the property of reductants but associated with the intrinsic characteristics of oxidation reactions involving MnO₄⁻ under acidic conditions. Fig. 1 also reveals that the loss of phenol follows the pseudo-first-order kinetics with MnO₄⁻ in 10-fold excess at pH 6.5–9.0, although the degradation rate of phenol changes with pH. The oxidation kinetics of aniline at pH 8.0 (Fig. S2†) and triclosan at pH 7.0 ³¹ further demonstrated that the decomposition of organic contaminants by MnO₄⁻ followed pseudo-first-order rate law with MnO₄⁻ in 10-fold excess under neutral and near-neutral conditions. The different oxidation kinetic behaviors over different pH ranges indicated different mechanisms of MnO₄⁻ oxidation.

Fig. 1 Oxidation kinetics of phenol by KMnO₄ over the pH range of 2.0–9.0. Experimental conditions: [phenol]₀ = 5.0 μM, [KMnO₄]₀ = 50 μM.

The gradual increase of the contaminant oxidation rate with time under acidic conditions indicated that the reduction product or intermediates of MnO₄⁻ accelerated contaminant oxidation. Jáky et al.⁵⁸ and Wiberg et al.⁵⁹ reported that the auto-acceleration of contaminant oxidation by MnO₄⁻ was ascribed to the presence of H⁺. Undoubtedly, the concentration of H⁺ was a crucial factor affecting the occurrence of auto-catalysis since auto-catalysis was only observed at pH 2.0–6.0, as shown in Fig. 1. Moreover, the rate of phenol oxidation increased with decreasing pH over the pH range of pH 2.0–6.0, indicating that phenol oxidation by permanganate was favored by H⁺. However, if the reaction was catalyzed by H⁺, the auto-acceleration of phenol and aniline oxidation by MnO₄⁻ would not be observed science the H⁺ concentration was kept constant during reaction. Therefore, H⁺ might mainly influence the generation and properties of the reduction products of MnO₄⁻ which resulted in the auto-acceleration of contaminant oxidation.

To identify the species responsible for the auto-acceleration, the change of the UV-vis spectra at 250–800 nm (where phenol and its oxidation products absorbed negligibly) during the course of the reaction between MnO₄⁻ and phenol over the pH range of 2.0–9.0 was monitored, as shown in Fig. 2. Under acidic conditions, the absorbance at 418 nm gradually increased accompanying with the oxidation of phenol by MnO₄⁻, indicating the generation of MnO₂.⁶⁰ To clarify the role of MnO₂ in the process of MnO₄⁻ oxidation, the experiments on phenol oxidation by colloidal MnO₂ under different conditions were conducted. As shown in Fig. 3, phenol could be oxidized by colloidal MnO₂ with greater rates than by MnO₄⁻ at pH 3.0 and 5.0, suggesting the highly oxidative activity of MnO₂ under this condition. Therefore, the auto-accelerated behavior of phenol oxidation by MnO₄⁻ at pH 2.0–6.0 should be mainly ascribed to the generation and accumulation of MnO₂ in this process, which was more reactive than MnO₄⁻ and would lead to an increasing phenol oxidation rate with time. However, the reactivity of colloidal MnO₂ dropped considerably with increasing pH and was much lower than MnO₄⁻ at pH > 6.0, as shown in Fig. 1 and 3C, which accounted for the disappearance of auto-acceleration behavior under near-neutral conditions.

Fig. 2 Variation of the UV-vis spectra during the course of the reaction between KMnO₄ and phenol at different pH. Experimental conditions: [phenol]₀ = 5.0 μM, [KMnO₄]₀ = 50 μM.

Fig. 3 Oxidation kinetics of phenol by colloidal MnO₂ (A) at pH 3.0, (B) at pH 5.0, (C) at pH 8.0. Experimental conditions: [phenol]₀ = 5.0 μM, [MnO₂]₀ = 50 μM.

To gain additional insight into the role of MnO₂ in accelerating contaminant oxidation by MnO₄⁻, the kinetics of phenol oxidation by MnO₄⁻ in the presence of excessive MnO₂ was determined. As shown in Fig. 4, although the dosing of colloidal MnO₂ greatly accelerated phenol degradation by MnO₄⁻ at pH 5.0, the auto-acceleration of phenol oxidation in this process was not observed. The enhanced phenol oxidation by MnO₄⁻ should be mainly ascribed to the oxidative ability of colloidal MnO₂ under this condition. However, the added MnO₂ promoted the degradation of phenol from the very beginning and masked the contribution of MnO₂ in situ generated from MnO₄⁻ reduction. Therefore, it could be concluded that it was the accumulation of in situ formed MnO₂ leading to the auto-acceleration of phenol oxidation by MnO₄⁻ at pH 2.0–6.0.

Fig. 4 Oxidation kinetics of phenol by KMnO₄ + colloidal MnO₂ at pH 5.0. Experimental conditions: [phenol]₀ = 5.0 μM, [MnO₂]₀ = 50 μM, [KMnO₄]₀ = 50 μM.

Based on the proposed degradation mechanisms of phenol by permanganate under acidic conditions, the degradation rate of phenol can be expressed as follows:

where k₁ and k₂ are the second-order rate constants of phenol oxidation by permanganate and in situ formed manganese dioxide, respectively. [MnO₄⁻]₀ is the initial concentration of MnO₄⁻, which was assumed to be constant during the reaction because it is dosed in large excess. [MnO₂]_t is the concentration of MnO₂ at time t. The concentration of MnO₂ increases with the proceeding of the reaction, resulting in the gradual increase in the oxidation rate of phenol under acidic condition. The excellent correlation of MnO₂ concentration with the increase in the phenol oxidation rate, as shown in Fig. S3,† verifies the acceleration of phenol oxidation by the in situ generated MnO₂.

3.2. The auto-inhibiting phenomenon in phenol oxidation by MnO₂

As revealed by Fig. 3, increasing pH from 3.0 to 8.0 resulted in a great decrease in phenol oxidation by MnO₂, which should be mainly ascribed to the drop in the redox potential of MnO₂/Mn²⁺ with increasing pH and the decrease in the adsorption of phenol on colloidal MnO₂.⁶¹ Another obvious trend demonstrated by Fig. 3 is that the removal rate of phenol by MnO₂ decreased with time over the pH range of 3.0–8.0. Fig. S4† shows that the absorbance of colloidal MnO₂ at 418 nm dropped negligibly, indicating that the auto-inhibiting behavior of phenol oxidation by MnO₂ was not associated with the large consumption of MnO₂. Previous studies demonstrated that cations could be absorbed by MnO₂ and influence the surface characteristics and oxidation properties of MnO₂.^38,62 Mn(II), the product of MnO₂ reduction, might be absorbed on the colloidal MnO₂ and inhibit the oxidation of phenol by MnO₂, which is a surface-mediated reaction. Therefore, the influence of Mn(II) on the oxidation kinetics of phenol by colloidal MnO₂ was examined at pH 3.0 and shown in Fig. 5. As expected, the oxidation of phenol by colloidal MnO₂ was retarded by dosing Mn(II) before the initiation of reaction. Furthermore, the auto-inhibiting phenomenon in the process of phenol oxidation by MnO₂ disappeared in the presence of supplementary Mn(II). This should be ascribed to the fact that the amount of Mn(II) dosed to the solution was much greater than that generated during reaction and thus masked the inhibiting effect of Mn(II) generated during phenol oxidation by MnO₂. In short, Mn(II) negatively affected the oxidative activity of MnO₂ because of its adsorption of on the surface of MnO₂ and occupation of the surface sites of MnO₂.

Fig. 5 Influence of Mn(II) on the oxidation kinetics of phenol by colloidal MnO₂ at pH 3.0. Experimental conditions: [phenol]₀ = 5.0 μM, [MnO₂]₀ = 50 μM.

3.3. Auto-accelerating of phenol oxidation by MnO₂ in the presence of PP

Surprisingly, the auto-inhibiting trend of phenol oxidation by MnO₂ disappeared but an auto-accelerating behavior was observed in the presence of PP at pH 3.0, as illustrated in Fig. 6A. PP was known as the ligands which could form strong complexes with Mn(III) (Mn(III)–PP) and stabilize Mn(III).^53,63 The conditional stability constants for Mn(III) complexes reported in the literature was listed as follows,⁶³

Mn(III) + PP ⇔ Mn(III)–PP, K = 1031.35 (at pH 8.0) (2)

Fig. 6 (A) Oxidation kinetics of phenol by colloidal MnO₂ in the presence of PP at pH 3.0. (B) Absorbance at 418 nm vs. absorbance at 258 nm for the reduction of MnO₂ by phenol in the presence of PP at pH 3.0. Experimental conditions: [phenol]₀ = 5.0 μM, [MnO₂]₀ = 50 μM, [PP]₀ = 5 mM.

Mn(III) stabilized by PP, Mn(III)–PP, has an absorbance peak at 258 nm, which was commonly used in studies on the role of Mn(III) in various processes.⁶³ Fig. S5† shows the change of the absorbance at 418 nm, the characteristic absorbance of colloidal MnO₂, and 258 nm with reaction time during phenol oxidation by MnO₂ in the presence of PP. Compared to the case without PP (Fig. S4†), much more MnO₂ was reduced due to the presence of PP. To determine the balance of manganese species, the linear relationship between the absorbance at 418 nm and 258 nm was constructed as follows.

Since the Mn(III)–PP complex absorbs negligibly at 418 nm,⁶⁰ one can express the absorbance at 418 nm as:

where is the molar absorptivity of MnO₂ at 418 nm and c_t is the concentration of MnO₂ at time t.

Since both MnO₂ and Mn(III)–PP absorb at 258 nm, the absorbance at this wavelength is given by eqn (4):

where and are the molar absorptivities of MnO₂ and Mn(III)–PP at 258 nm, respectively. c₀ is the initial MnO₂ concentration. Based on eqn (3) and (4), it is easy to deduce:

The excellent linearity of A₄₁₈ with A₂₅₈, as shown in Fig. 6B, suggesting the negligible accumulation of Mn(II) during phenol oxidation by MnO₂ in the presence of PP. The oxidation rate of phenol by MnO₂ in the presence of PP did not drop with time (Fig. 6A), further confirming the negligible accumulation of Mn(II). Moreover, the dosing of PP greatly enhanced phenol oxidation by MnO₂, which should be associated with the high reactivity of Mn(III)–PP, as demonstrated in Fig. 7. Therefore, the presence of PP changes the manganese transformation route during phenol oxidation by MnO₂. Without the presence of PP, MnO₂ was reduced to Mn(II) ions, which tend to decrease the reactivity of MnO₂. However, in the presence of excessive PP, MnO₂ was reduced to Mn(III)–PP. Mn(III)–PP was much more reactive than colloidal MnO₂ at pH 3.0 (Fig. 3A and 7), resulting in an increase in the oxidation rate of phenol by colloidal MnO₂ with time in the presence of excessive PP.

Fig. 7 Oxidation kinetics of phenol by Mn(III)–PP complex at pH 3.0. Experimental conditions: [phenol]₀ = 5.0 μM, [Mn(III)–PP]₀ = 50 μM, [PP]₀ = 5 mM.

Similarly to phenol oxidation by MnO₂ (Fig. 3), a gradual decrease in the rate of phenol oxidation by Mn(III)–PP was observed (Fig. 7). Compared to MnO₂ and MnO₄⁻, more Mn(III)–PP was needed to oxidize the same amount of contaminants as one Mn(III)–PP molecule could only transfer one electron. Fig. S6† showed the significant decrease of Mn(III)–PP concentration during phenol oxidation, resulting in the decrease of oxidation rate of phenol with time by Mn(III)–PP.

Based on the above discussion, the oxidation rate of phenol by manganese dioxide in the presence of PP under acidic conditions could be expressed as follows:

where k₃ is the second-order rate constant of phenol oxidation by in situ formed Mn(III)–PP. The excellent correlation demonstrated in Fig. S7† further confirmed the rationality of the proposed mechanisms that the acceleration of phenol oxidation by manganese dioxide in the presence of PP under acidic conditions was mainly attributed to the generation of Mn(III)–PP.

3.4. Auto-acceleration of phenol oxidation by MnO₄⁻ in the presence of PP

The influence of PP on the oxidation of phenol by MnO₄⁻ at pH 3.0 is shown in Fig. 8A. Obviously, the decomposition of phenol displayed an auto-acceleration behavior since the reaction rate increased as the reaction progressed. Compared with the case without PP (Fig. 1), PP enhanced phenol oxidation significantly and shortened the reaction time over 10 times to achieve the removal of 99% phenol. The variation of UV absorbance at 525 nm (arising from MnO₄⁻), 418 nm (arising from MnO₂), and 258 nm (arising from both MnO₂ and Mn(III)–PP) with reaction time was monitored, as shown in Fig. S8.† Fig. S8† indicated that the concentration of MnO₄⁻ decreased progressively while negligible MnO₂ was generated since the absorbance at 418 nm was always very low throughout the reaction. However, the concentration of Mn(III)–PP increased during the reaction. Based on the consumption of MnO₄⁻ and the formation of Mn(III)–PP, the balance of manganese species was constructed following the methods of deriving eqn (2)–(4) and described as follows,where and are the molar absorptivities of MnO₄⁻ and Mn(III)–PP at 258 nm, respectively. is the molar absorptivity of MnO₄⁻ at 525 nm, c₀ is the initial MnO₄⁻ concentration, and c_t is the concentration of MnO₄⁻ at time t. The excellent correlation of the absorbance at 525 nm and 258 nm (Fig. 8B) indicated that manganese mainly existed as MnO₄⁻ and Mn(III)–PP in the solution. Therefore, the higher activity of Mn(III)–PP towards phenol compared to MnO₄⁻ resulted in the auto-acceleration of phenol oxidation by MnO₄⁻ in the presence of excessive PP.

Fig. 8 (A) Oxidation kinetics of phenol by KMnO₄ in the presence of PP at pH 3.0. (B) Absorbance at 525 nm vs. absorbance at 258 nm for the reduction of permanganate by phenol in the presence of PP at pH 3.0. Experimental conditions: [phenol]₀ = 5.0 μM, [KMnO₄]₀ = 50 μM, [PP]₀ = 5 mM.

According to the above discussion, the oxidation rate of phenol by permanganate in the presence of PP under acidic conditions could be expressed as follows:

There is an excellent correlation between the concentration of generated Mn(III)–PP and the enhancement of phenol oxidation rate, as illustrated in Fig. S9,† verifying the rationality of the proposed mechanisms that the in situ formed Mn(III)–PP was primarily responsible for the acceleration of phenol oxidation by permanganate in the presence of PP under acidic conditions. However, Mn(III) was quite unstable and easily disproportionated to Mn(II) and MnO₂.⁶⁴ Thus, the accumulation of Mn(III) was impossible to be responsible for the auto-acceleration of phenol oxidation by MnO₄⁻ in the absence of PP or other ligands. The only candidate for reactive intermediates that accounted for the auto-acceleration of contaminant oxidation by MnO₄⁻ in the absence of ligands is MnO₂.

4. Conclusions

In this paper, the kinetic law of phenol oxidation by MnO₄⁻ and MnO₂ was investigated in the presence and absence of PP under acidic conditions. Auto-acceleration was observed during phenol oxidation by MnO₄⁻ under acidic conditions due to the generation and accumulation of highly active MnO₂. Conversely, an auto-inhibiting trend appeared in the oxidation of phenol by MnO₂, resulting from the generation of Mn(II) and its subsequent absorption on the surface of colloidal MnO₂. The oxidation of phenol by MnO₄⁻ and MnO₂ was greatly enhanced in the presence of PP and the oxidation rate increased with time. The auto-accelerating phenomena in the process of phenol oxidation by MnO₄⁻ or MnO₂ in the presence of excessive PP were mainly attributed to the generation and accumulation of Mn(III)–PP, which was much more reactive than MnO₄⁻ and MnO₂ under acidic conditions. Considering the influence of in situ formed manganese intermediates on the degradation of contaminants by manganese oxides, it is a promising method to enhance the degradation of contaminants by promoting the generation of highly active manganese intermediates in water treatment.

Acknowledgements

This work was supported by the National Natural Science Foundation of China (Grant 21522704), the Major Science and Technology Program for Water Pollution Control and Treatment (2012ZX07403-001), and the Fundamental Research Funds for the Central Universities.

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Footnote

† Electronic supplementary information (ESI) available. See DOI: 10.1039/c6ra10196h

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