Role of inorganic ions and dissolved natural organic matters on persulfate oxidation of acid orange 7 with zero-valent iron

Abstract

The impacts of common anions and organic matter, initial pH and PS dosage on the oxidation of acid orange 7 (AO7) by persulfate (PS) activated with zero-valent iron (ZVI) were investigated. The present findings revealed that maximum AO7 decolorization occurred at pH 3.0, and increasing the system pH resulted in a greater decrease in AO7 decolorization rates. AO7 decolorization efficiency was 100% at 120 min when the molar ratio of PS : AO7 was 5 : 1. Interestingly, ClO₄⁻, CH₃COO⁻ and humic acid (HA) were found to accelerate AO7 decolorization rates while other anions retarded AO7 decolorization in the following sequence: NO₂⁻ > H₂PO₄⁻ > HPO₄²⁻ > EDTA > SO₄²⁻ > CO₃²⁻ > HCO₃⁻ > NO₃⁻ > Cl⁻. ClO₄⁻, CH₃COO⁻ and HA with 50 mM, 10 mM and 1.0 mg L⁻¹, respectively were found to be the optimal concentrations for AO7 decolorization. The removal efficiencies of AO7 were decreased by 90.3%, 51.5%, 58.7% and 38.2%, respectively over 120 min in addition of NO₂⁻ (50 mM), H₂PO₄⁻ (50 mM), HPO₄²⁻ (50 mM) and EDTA (50 mM). The other anions including SO₄²⁻, CO₃²⁻, HCO₃⁻, NO₃⁻ and Cl⁻ led to a decrease change of less than 20%. The mechanisms for the influence were complexation reactions with Fe²⁺ generated from ZVI, consuming of sulfate radicals (SO₄⁻˙) by scavenging reactions, and oxidation reactions involving inorganic ions. The reason for the acceleration by CH₃COO⁻ and HA was probably through acting as an electron ‘shuttle’ and facilitating electron transfer from the ZVI surface to PS and resulted in more Fe²⁺ and SO₄⁻˙. However, the acceleration caused by ClO₄⁻ was presumably ascribed to the oxidizing of ZVI directly by ClO₄⁻ to produce more Fe²⁺.

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1. Introduction

Synthetic dyes have received increasing attention in recent years due to their wide use in many industries such as textile, cosmetic, pulp and paper.¹ As one of the most widely used synthetic dyes, azo dyes are ubiquitous in environments² because of their high solubility in water, which would cause harm to aquatic organisms and human health due to their toxic and potential carcinogenic nature.¹ Another potential disadvantage is from the fact that azo dyes are difficult to degrade by biological treatment methods due to their complex structure and stability.³

Advanced oxidation processes (AOPs) involve the generation of free radicals, notably hydroxyl radicals (HO˙) and sulfate radicals (SO₄⁻˙) that are highly oxidative and capable of degrading a wide range of organic compounds.⁴ Typical AOP was based on the generation of HO·, such as the Fenton reaction.⁵ Recently, AOP that generates nonselective SO₄⁻˙ by activation of persulfate ion (S₂O₈²⁻, PS)⁶ or peroxymonosulfate (PMS)⁷ has attracted a great deal of interest. Compared to the Fenton reagent and other oxidants, some properties of PS including high aqueous solubility, more chemically stable in subsurface, relatively low cost, easy storage and transport make it to be a promising oxidant of in situ chemical oxidation (ISCO) which is a technique used to remediate contaminated soil and groundwater systems.⁶ SO₄⁻˙ has a high oxidizing potential of 2.5–3.1 versus normal hydrogen electrode (NHE),⁶ which makes it an excellent oxidant for degrading a wide range of recalcitrant and/or toxic organic pollutants in water and soil.^8–10 Generally, SO₄⁻˙ can be produced from the activation of PS by UV-Vis,^8,9 heat,¹⁰ and transition metal ions.^11–13

Photochemical activation: S₂O₈²⁻ + hν⁻ → 2SO₄⁻˙ (1)

Thermal activation: S₂O₈²⁻ + heat → 2SO₄⁻˙ (2)

Chemical activation: S₂O₈²⁻ + Mⁿ⁺ → SO₄⁻˙ + M⁽ⁿ⁺¹⁾⁺ + SO₄²⁻ (3)

Among the various activation methods, zero-valent iron (ZVI) activation of PS has appeared to be a promising method.^14–16 ZVI activation is a cost-effective, efficient and environmentally friendly technology.^14,17 Additionally, ZVI not only serves as a slow-releasing source of dissolved Fe²⁺, but also avoids adding other anions that can lead to scavenging of SO₄⁻˙ and possibly reducing the oxidation efficiency.^18,19 Because the competing side reactions of SO₄⁻˙ with these anions rather than the target pollutant would occur.

Recently, one of the limitation factors in the use of PS for degradation of organic pollutants is the reactions between the produced radicals and nontarget chemical species that are naturally occurring or anthropogenic present in the wastewater.^19,20 Besides the refractory contaminants, waster waters usually contain a certain amount of other substances such as inorganic ions and common dissolved natural organic matters. It was reported that the concentrations levels of sulfate and chloride anions in groundwater vary from 0.1 to 100 mM.²¹ Other naturally abundant inorganic anions such as nitrate, carbonate and phosphate anions are frequently present in water, wastewater or seawater with various concentrations. The existence of these substances may affect the degradation rate of target contaminant by serving as proton donors, electron shuttles, or competing for electrons, and thereby affecting the oxidation efficiency of the target pollutant.²² Also, common organic matters and anions have the potential to impact pathway and kinetic of oxidation reactions both as radical scavengers and metal complexing agents.^21,23 This may result in limiting the reactivity of SO₄⁻˙ by the presence of background ions in wastewaters. Based on Liang's research, the Cl⁻ would exhibit an inhibition effect on the degradation of TCE by thermal/PS once the concentration of Cl⁻ was greater than 0.2 mM.²⁴ Carbonates also reduced the decomposition rates of contaminant and oxidant in the PS oxidation system.²⁵ While effects of anions on heat, UV light and ferrous ion activation methods have been studied,^26,27 literature of anions on the behavior of ZVI activating of PS is very limited. Recently, we have reported influence of particle size of ZVI and dissolved silica on the reactivity of activating PS for degradation of AO7.²⁶ However, the effects of other common organic matters and anions on the decolorization of AO7 are also needed to understand. Moreover, in the research of influence of anions in the PS systems, researchers did not deeply study the retardation or promoting mechanisms. Therefore, in the present study we aim to better understand the behavior of the common anions and organic matters on the SO₄⁻˙-based treatment of organic contaminants, using PS/ZVI/AO7 as a model AOP treatment technology. Therefore, the objective of this research is to gain insight into: (1) the influence of both acidity anions and alkalinity anions on the removal of AO7 in the PS/ZVI system, (2) the impacts of dissolved natural organic matters on the decomposition of AO7 and PS, (3) the influence mechanism of these common organic matters and anions.

2. Materials and methods

2.1 Materials

The ZVI (purity > 99%, approx. 150 μm), humic acid (HA, fulvic acid (FA) ≥ 90%), and sodium perchlorate (NaClO₄, 99.0%) were obtained from Aladdin chemistry Co., Ltd (Shanghai, China). The total surface area (α_s) of ZVI was 2.1518 m² g⁻¹ according to the N₂ isothermal adsorption. AO7 (purity > 99.0%) was purchased from Tokyo Chemical Industry (Japan). Sodium persulfate (PS, Na₂S₂O₈, 98.0%), sodium chloride (NaCl, 99.5%), sodium sulfate (Na₂SO₄, 99.0%), sodium bicarbonate (NaHCO₃, 99.5%), sodium carbonate (Na₂CO₃, 99.0%), sodium phosphate dibasic trihydrate (Na₂HPO₄·3H₂O, 99.0%), sodium nitrate (NaNO₃, 99.0%), sodium nitrite (NaNO₂, 99.0%), ethylene diamine tetraacetic acid disodium (EDTA-2Na, 98.0%), potassium iodide (KI, 99.0%), and other chemicals used were purchased from Sinopharm Chemical Reagent Co., Ltd (Beijing, China). All chemicals were used as received without any further purification.

2.2 Experimental procedure

The experimental setup was similar to our previously study.¹⁵ Reactions were carried out in a 250 mL Erlenmeyerl flask at 25 ± 0.2 °C. Fifty millilitres of the prepared AO7 (0.4 mM) and PS (4 mM) stock solutions were added simultaneously to the reactor, giving initial concentrations of AO7 and PS of 0.2 mM and 2.0 mM, respectively. Reaction mixtures in the flask were subsequently added 0.5 mL of inorganic ions or organic matters stock solution with high concentration. The flask was open to the atmosphere and shaken at 180 rpm in a rotary shaker (ZHWY-20102C, Shanghai, China). All reactions were initiated by adding ZVI, then samples were withdrawn in the predetermined time intervals, and then a certain amount of ethanol (1.0 mL of ethanol for each 1.0 mL sample) was added to quench the reaction.²⁸ The supernatant was filtered through a 0.45 μm membrane filter and analyzed for AO7, PS, Fe²⁺ and total dissolved iron. In order to avoid potential complications by buffers or ionic strength, all experiments except the effect of initial pH were performed without pH adjustment and to give initial pHs of 3.8. But the pHs varied dramatically after adding inorganic anions or organic matters, the pH of reaction solutions containing inorganic anions or organic matters was displayed in Tables S1 and S2.† In the experiment of effect of initial pH, the initial pH value was adjusted with 1.0 M sodium hydroxide (NaOH) or sulfuric acid (H₂SO₄).

2.3 Analytical methods

The absorbance of Acid Orange 7 was measured in visible spectra at the characteristic wavelength of AO7 (λ_max = 486 nm) using a UV-Vis spectrophotometer (UV2301 II, Shanghai, China). Decolorization efficiency (η) was calculated based on the following equation: η (%) = (A₀ − A)/A₀ × 100, where the A₀ and A were the absorbance of the sample at time 0 and t, respectively. The PS residual was determined by spectrophotometric determination with potassium iodide,²⁹ and the concentrations of ferrous iron and dissolved iron were measured with 1,10-phenanthroline at a wavelength of 510 nm.³⁰ The concentrations of some anions were measured by ion chromatography (DX 500, USA). The pH was monitored by pH meter (Shanghai LeiCi PHS-25) equipped with a pH electrode.

3. Results and discussion

3.1 Effect of initial pH and PS dosage

The pH value of aqueous solution plays an important role in the degradation of organic compounds in the advance oxidation processes. To confirm the effect of initial pH on the decolorization rate of AO7, a series of experiments were carried out with initial pH values ranging from 3.0 to 11.0 (seen in Fig. 1). Clearly, AO7 could completely decolor over a wide pH range of 3.0–9.0. Moreover, the solution pH significantly influenced the AO7 decolorization rates, which decreased with the increase of the initial pH values. It can be derived from this result that the acidic solution was more favorable than neutral and alkaline solutions for generating of SO₄⁻˙. The AO7 decolorization under different initial pH was observed to fit the pseudo-first order kinetic model well ([AO7]/[AO7]₀ = exp(−k_obst)) with a correlation coefficient R² value greater than 0.96, where the k_obs represented the pseudo-first-order rate constant. It can be seen (seen in Fig. 1b) that the rate constants decreased significantly when the pH increased from 3.0 to 5.0. However, the AO7 decolorization rate constants showed small extent decrease while the solution pH values increased from 5.0 to 9.0.

Fig. 1 (a) The decolorization of AO7 under different initial pH values; (b) pseudo-first-order rate constants of AO7 versus initial pH values. Experiment condition: PS = 2 mM; AO7 = 0.2 mM; ZVI = 0.5 g L⁻¹; T = 25 °C.

PS can be decomposed to produce SO₄⁻˙ with the activation of catalysts, which makes a dominant contribution to removing organic contaminants. The influence of the molar ratios of PS to AO7 on the AO7 decolorization was studied. As illustrated in the Fig. 2, the concentrations of PS appeared to be one of the crucial factors affecting the AO7 decolorization. The AO7 decolorization efficiency was noticeable higher at the molar ratio (PS : AO7) of 5 : 1 than that of 1 : 1. AO7 was decolorized completely in 120 min when the molar ratio of PS : AO7 exceed 5 : 1.

Fig. 2 The effect of PS dosage on the decolorization of AO7 versus reaction time. Experiment condition: AO7 = 0.2 mM; ZVI = 0.5 g L⁻¹; initial pH = 3.8 ± 0.1; T = 25 °C.

3.2 The role of common acidity anions

Usually, a great amount of dissolved inorganic ions may be present initially in the wastewater or formed as end products from the compounds undergoing degradation. Therefore, in order to clear the role of common acidity anions on the degradation of organic pollutants in the ZVI/PS system, the decolorization rate of AO7 was measured in the presence of acidity anions with the concentration of 50 mM. As shown in Fig. 3, it can be obviously found that the addition of different acidity inorganic ions led to different impact on decolorization of AO7 with ZVI activation of PS.

Fig. 3 (a) The effect of common acidity anions on the decolorization of AO7; (b) the effect of common acidity anions on the decomposition of PS. Experiment condition: PS = 2 mM; AO7 = 0.2 mM; ZVI = 0.5 g L⁻¹; Cl⁻ = 50 mM; ClO₄⁻ = 50 mM; SO₄²⁻ = 50 mM; NO₃⁻ = 50 mM; NO₂⁻ = 50 mM; T = 25 °C.

3.2.1 Nitrite and nitrate ions. Nitrite ion (NO₂⁻) exhibited relatively deceleration on the decolorization of AO7, only about 10% of the AO7 was removed from solution over 120 min while the AO7 decolorization rate was 100% without NO₂⁻ under the identical conditions. However, it was interesting to note that the addition of nitrate ion (NO₃⁻) showed a much smaller extent deceleration on the AO7 decolorization, whereas the AO7 decolorization efficiency were 90.0% and 94.3% within 60 min in the presence and absence of NO₃⁻, respectively. In fact, ZVI has been known to reduce nitrate to nitrite, nitrogen, ammonium and ammonia depending upon experimental conditions, as well as zero valent aluminum and zero valent magnesium.³¹ Moreover, NO₃⁻ does not complex with Fe³⁺ or Fe²⁺ measurably, nor does it react with hydroxide radical.³² From which we can infer that the little inhibitory effect of NO₃⁻ may be ascribed to that NO₃⁻ competed with PS on electrons generated from ZVI corrosion. The deceleration decolorization caused by NO₃⁻ also confirmed the hypothesized mechanism of heterogeneous activation of PS, involving direct electron transfer from ZVI (see eqn (4)).

Fe⁰ + 2S₂O₈²⁻ → Fe²⁺ + 2SO₄²⁻ + 2SO₄⁻˙ (4)

The remarkable effect of NO₂⁻ on the decolorization of AO7 was due to the transformation of NO₂⁻ to NO₃⁻ (see eqn (5)), which consumed large amounts of SO₄⁻˙ and caused the final pH lower than initial pH.

2SO₄⁻˙ + H₂O + NO₂⁻ → NO₃⁻ + 2SO₄²⁻ + 2H⁺ (5)

For NO₃⁻ was detected and increased when the NO₂⁻ concentration decreased during the reaction (data not shown). On the other hand, although the PS decomposition was also retarded in the presence of NO₂⁻ (Fig. 3b), the retardation effect was smaller than that of AO7. Based on these results, we speculated that the scavenging of SO₄⁻˙ by NO₂⁻ was not mainly resulted in the deceleration effect. It was likely, however, the retardation effect on the decomposition of PS, which caused the low concentration of SO₄⁻˙. Further experiments were conducted in various concentrations of NO₂⁻. As shown in Table S3,† the pseudo-first-order rate constants (k_obs) of the decolorization of AO7 in contact with NO₂⁻ concentrations ranging from 0 to 100.0 mM were obtained. The oxidation of AO7 exhibited great decrease in the rate constants from 0.048 min⁻¹ to 0.001 min⁻¹ with NO₂⁻ concentration increase from 0 to 50.0 mM. However, when the concentration of NO₂⁻ exceeded 50.0 mM, the k_obs increased to 0.002 min⁻¹ and the half-life decreased correspondingly.

3.2.2 Sulfate ion. Evidently, the retention on the oxidation of AO7 was observed when the sulfate ion (SO₄²⁻) was added into the PS/ZVI system (seen in Fig. 3a) even though they were poor free radical scavengers.³³ This effect was also reflected in the PS decomposition. However, compared to the control system without any anion, the presence of SO₄²⁻ led to a decrease of 52% change in decomposition of PS (seen in Fig. 3b) and only 14% change in decolorization of AO7 (seen in Fig. 3a). This verified that the retardation effect on decolorization of AO7 was not ascribed to scavenging of SO₄⁻˙ by SO₄²⁻. In contrast, the appearance of SO₄²⁻ made SO₄⁻˙ more efficiently towards oxidation of AO7. This result can be attributed to the fact that SO₄²⁻ underwent complex reactions with Fe³⁺ and Fe²⁺ and formed a mixture of FeSO₄⁺ and Fe(SO₄)₂⁻ complexes, which decreased the concentration of Fe²⁺ and reduced Fe³⁺ through coordination.⁹ In addition, it has been reported³⁴ that the presence of SO₄²⁻ could reduce oxygenation of Fe²⁺ in neutral and slightly acidic solution because of the formation of ion pairs (FeSO₄) that are more difficult to oxidize. Therefore, it was likely that the retardation effect caused by SO₄²⁻ could be overcome by extending the reaction time in this system.

3.2.3 Chloride and perchlorate ions. Based on research results of Liang et al.,²³ there was hardly any interference emerging in the TCE degradation with chloride levels below 0.2 M on PS oxidation of TCE at 20 °C. Another study²⁰ showed that iron activation at neutral pH was not affected significantly by Cl⁻ with concentrations of 5.0 and 50.0 mM. In this study, the Cl⁻ and ClO₄⁻ revealed little effect both on the AO7 decolorization and PS decomposition (Fig. 3a). Compared to the AO7 decolorization, the presence of Cl⁻ resulted in a greater degree of retardation on the PS decomposition. These findings illustrated that the deceleration was attributed to the effective scavenging of SO₄⁻˙ through the reactions (see eqn (6) and (7))³⁵ of Cl⁻ with SO₄⁻˙ generated from PS activated by ZVI.

SO₄⁻˙ + Cl⁻ → SO₄²⁻ + Cl˙ (6)

Cl˙ + Cl⁻ → Cl₂⁻˙ (7)

Perchlorate ions (ClO₄⁻) are similar to NO₃⁻, they do not react with Fe³⁺ or Fe²⁺ through complex reaction, nor do they react with free radicals. In the presence of ClO₄⁻, Fe(II) exists as Fe²⁺ and Fe(OH)⁺ at pH < 8.³⁶ Hence, the reaction between Fe²⁺ and PS was not suppressed. However, it was interesting to find that pseudo-first-order rate constants for PS oxidation of AO7 increased from 0.048 min⁻¹ to 0.074 min⁻¹ (seen in Table S4†) when the ClO₄⁻ concentration increased from 0 to 50.0 mM. This indicated that the AO7 decolorization was drastically enhanced by ClO₄⁻, which was due to a probable reaction between ZVI and ClO₄⁻ (see eqn (8)). The final pH was a little higher than the initial pH of PS + AO7 + ZVI + ClO₄Na system also supported this conclusion. The final pH was usually lower than initial pH in SO₄⁻˙-based AOPs.^14,15 Once the ClO₄⁻ concentrations exceed 50.0 mM, the rate constants decreased but they were still higher than the rate constant without ClO₄⁻.

ClO₄⁻ + 4Fe⁰ + 8H⁺ → 4Fe²⁺ + Cl⁻ + 4H₂O (8)

3.3 The role of common alkalinity anions

As evidence, the negativity of common alkalinity anions was greater than that of common acidity anions on the decolorization efficiency of AO7, which was showed in the Fig. 4.

Fig. 4 (a) The effect of alkalinity anions on the decolorization of AO7 versus reaction time; (b) the effect of alkalinity anions on the decomposition of PS versus reaction time. Experiment condition: PS = 2 mM; AO7 = 0.2 mM; ZVI = 0.5 g L⁻¹; CO₃²⁻ = 50 mM; HCO₃⁻ = 50 mM; HPO₄²⁻ = 50 mM; H₂PO₄⁻ = 50 mM; CH₃COO⁻ = 50 mM; T = 25 °C.

3.3.1 Carbonates and bicarbonate ions. Carbonates ion (CO₃²⁻) and bicarbonate ion (HCO₃⁻) are well known buffer ions and often adopted to adjust the pH values, implying that they are expected to be extremely important species. It was reported that addition of CO₃²⁻ and HCO₃⁻ increased oxidation of Fe²⁺ forms dramatically, which was believed that this effect was due to the higher reactivity of FeCO₃ (than Fe²⁺ or FeOH⁺) towards H₂O₂.³⁷ Furthermore, CO₃²⁻ are also believed to adsorb and inactivate catalytic and scavenging sites such as iron oxides.³⁸ As displayed in Fig. 4a, the PS reaction was extremely sensitive to CO₃²⁻ remaining in the solution, but the addition of HCO₃⁻ inhibited the AO7 decolorization to a much smaller extent. After 120 min of the reaction with PS, AO7 was decolorized in 96.2% in the presence of HCO₃⁻. When CO₃²⁻ was introduced into the reaction solution, the amount of decolorized AO7 fell to 89.1%. The greater scavenging capacity of CO₃²⁻ > HCO₃⁻ was considered to be the major reason responsible for this phenomenon. The rate constant for the reaction of SO₄⁻˙ with HCO₃⁻ is about 4 times lower than that of SO₄⁻˙ with CO₃²⁻ (see eqn (9) and (10)). Moreover, the redox potential of CO₃⁻˙ is lower than that of HCO₃·.³⁹ Accordingly, the inhibition impact on oxidation of AO7 by CO₃²⁻ was more pronounced than that of HCO₃⁻. The HCO₃˙ and CO₃⁻˙ generated by the reaction of SO₄⁻˙ with HCO₃⁻ and CO₃²⁻ were reported to yield redox potential,²³ which can possibly destroy AO7. But the generating rate and the redox potential of these two free radicals were significantly lower compared to SO₄⁻˙, this revealed that the contribution of HCO₃˙ and CO₃⁻˙ towards the oxidation of AO7 might be negligible.

SO₄⁻˙ + CO₃²⁻ → SO₄²⁻ + CO₃⁻˙ (9)

SO₄⁻˙ + HCO₃⁻ → SO₄²⁻ + HCO₃˙ (10)

3.3.2 Hydrogen phosphate and dihydrogen phosphate ions. The addition of hydrogen phosphate ion (HPO₄²⁻) and dihydrogen phosphate ion (H₂PO₄⁻) showed retardation effects on the decolorization of AO7, but the delay was more pronounced in the case of H₂PO₄⁻. The removal efficiencies of AO7 in the presence of H₂PO₄⁻ and HPO₄²⁻ were decreased by 51.5% and 58.7%, respectively. The markedly decreasing rates on decomposition of PS caused by H₂PO₄⁻ were resulted from the complex compounds of H₂PO₄⁻ with Fe²⁺ or Fe³⁺ (see eqn (11) and (12)). These phosphate complexes are quite insoluble in neutral or mildly acidic solution.

Fe²⁺ + H₂PO₄⁻ → FeH₂PO₄⁺ (11)

Fe³⁺ + H₂PO₄⁻ → FeH₂PO₄²⁺ (12)

Accordingly, precipitation of Fe(III) phosphate complexes presumably reduced the reactivity species of Fe²⁺ towards activating of PS. The inhibiting effect of H₂PO₄⁻ on the rates of conversion of AO7 not only depended on the complexation of ferrous ions, but also resulted from competition with AO7 for SO₄⁻˙ because H₂PO₄⁻ reacted with SO₄⁻˙ to generate inorganic radicals (see eqn (13)).

SO₄⁻˙ + H₂PO₄⁻ → H₂PO₄˙ + SO₄²⁻ (13)

H₂PO₄˙, one of the strong oxidant species, react with most of the organic solutes with a high second-order rate constants that range from 10⁶ to 10⁹ M⁻¹ s⁻¹, but they are still less reactive than SO₄⁻˙. In the case of HPO₄²⁻, it is well known that HPO₄²⁻ are efficient scavengers of HO·,³⁴ maybe scavengers of SO₄⁻˙ as well (see eqn (14)). This hypothesis can be verified by the

SO₄⁻˙ + HPO₄²⁻ → HPO₄⁻˙+ SO₄²⁻ (14)

decomposition of PS, which showed less extent inhibition effect than that of AO7. Apparently, the inhibition differences between these two phosphate ions were attributed to the phosphate complexes of iron.

3.3.3 Acetate ion. The effect of acetate ion (CH₃COO⁻) on the decolorization of AO7 in PS/ZVI process was investigated (Fig. 4). Surprisingly, the decolorization rate underwent significant enhancement in the presence of 50.0 mM CH₃COO⁻ as compared to the control system with no added CH₃COO⁻. In contrast, the concentration of remaining PS with CH₃COO⁻ was greater than that of the control system. This indicated that CH₃COO⁻˙ behaved differently from other common alkalinity anions, which not only increased the AO7 decolorization rate but also decreased the scavenging reactions that competed with AO7 for SO₄⁻˙. Further study on the various concentration of CH₃COO⁻ were conducted, data presented in Table S5.† It can be observed that increasing the concentrations of CH₃COO⁻ from 0 to 10.0 mM resulted in an increase in rate constants from 0.048 to 0.064 min⁻¹. However, further increase of the CH₃COO⁻ concentrations beyond 10.0 mM led to gradual decrease in the rate constants. The concentrations of Fe²⁺, total dissolved iron, and CH₃COO⁻ were monitored (Fig. 5) during the reaction in order to verify the influence mechanism. Observation of the concentration of Fe²⁺, total dissolved iron indicated that the presence of CH₃COO⁻ could promote the release of Fe²⁺, which was one of the main species that can activate PS to produce SO₄⁻˙. Therefore, it was speculated that CH₃COO⁻ acted as an electron shuttle to promote the electron transfer from ZVI, and generate Fe²⁺ in this oxidation process. However, the speed of Fe²⁺ release was dramatically different from that of PS/Fe²⁺ system where supplied abundant Fe²⁺ instantaneously and resulted in a great amount of Fe²⁺ inactivating and declining AO7 and PS decomposition rate eventually, as shown in Fig. S1.† The results demonstrated that the slow release of Fe²⁺ from ZVI was better than the direct addition of Fe²⁺. It should be noted that the concentrations of CH₃COO⁻ decreased from 50 to 47.3 mM within 120 min, it was likely due to the formation of^·CH₂COO⁻ radical by H-abstraction reaction (see eqn (15)). This also potentially explained why the final pH was much lower than the initial pH in the PS + AO7 + ZVI + CH₃COONa system.

SO₄⁻˙ + CH₃COO⁻ → CH₂COO⁻˙ + SO₄²⁻ + H⁺ (15)

Fig. 5 The variation of iron ions in the presence and absence of CH₃COO⁻ versus reaction time. Experiment condition: PS = 2 mM; AO7 = 0.2 mM; CH₃COO⁻ = 50 mM; ZVI = 0.5 g L⁻¹; T = 25 °C.

This scavenging reaction caused the AO7 decolorization rate constant decreasing when the CH₃COO⁻ concentrations increased. Another mechanism for the enhancement effect might be the complex reactions. The reaction solution was clear all the time in the presence of CH₃COO⁻ while the solution gradually became turbid in the absence of CH₃COO⁻ because of the precipitation of Fe(III). This phenomenon helped to get the conclusion of complex reactions in the absence of CH₃COO⁻.

3.4 The role of common organic matters

3.4.1 Ethylene diamine tetraacetic acid. As one of the most widely used chelating agents, ethylene diamine tetraacetic acid (EDTA) is routinely used to remove heavy metal ions from hard water or in industrial cleaning.⁴⁰ More recently, EDTA was adopted as chelating agents to improve the treatment efficiency by slow releasing Fe²⁺ in AOPs.^41,42 Therefore, the impact of EDTA on the degradation of organic pollutant seems to be important in AOPs. It was seen (seen in Fig. 6) that both AO7 and PS decomposition were retarded in the presence of 50 mM EDTA. The control system with PS but no ZVI, AO7 and PS were almost not decomposed, which revealed that EDTA showed little influence on the activation of PS. Consequently, this was speculated that retardation effect was due to chelating Fe²⁺ by EDTA. On the other hand, oxidation of EDTA by SO₄⁻˙ also resulted in AO7 decolorization decrease.

Fig. 6 (a) The effect of common organic matters on the decolorization of AO7; (b) the effect of common organic matters on the decomposition of PS. Experiment condition: PS = 2 mM; AO7 = 0.2 mM; ZVI = 0.5 g L⁻¹; EDTA = 50 mM; HA = 5.0 mg L⁻¹; T = 25 °C.

3.4.2 Humic acid. Humic acid (HA) is one of the important component of natural systems, the role of HA on the PS oxidation reaction has not been reported, to the best of our knowledge. The data (Fig. 6) showed a small acceleration on the AO7 decolorization in the presence of 5.0 mg L⁻¹ HA, the similar phenomenon was also displayed in the decomposition of PS. But AO7 and PS were not decomposed without ZVI, indicating that HA could not activate PS to produce SO₄⁻˙. When HA decreased to 0.5 mg L⁻¹ in the reaction solutions, the rate constants of AO7 decolorization increased from 0.048 to 0.101 min⁻¹ (Table S6†). However, HA showed a clearly deceleration on the AO7 decolorization when its concentration was greater than 5.0 mg L⁻¹. The rate constants in the presence of 1.0, 5.0, 7.5 and 10 mg L⁻¹ of HA were 0.105, 0.049, 0.041 and 0.045 min⁻¹, respectively. Based on the results, we speculated that HA acted as an electron shuttle and the chelation of iron made its influence on AO7 decomposition very interesting. The concentrations of Fe²⁺ and total dissolved iron also increased dramatically in the presence of HA (5.0 mg L⁻¹) as compared to the control system without HA (Fig. 7). It was reported that primary functional groups including quinones, carboxylic acids, alcohols, and ketones in humic substances acted as soluble electron carriers to facilitate the degradation of pollutants by accepting electrons from ZVI and ‘shuttling’ the electrons to the H₂O₂ in Fenton reactions.⁴³ Therefore, the presence of HA boosted the electron transfer from ZVI surface to the PS and resulted in the accelerating formation of Fe²⁺ and SO₄⁻˙ in the PS/ZVI system.

Fig. 7 The variation of iron ions in the presence and absence of HA versus reaction time. Experiment condition: PS = 2 mM; AO7 = 0.2 mM; ZVI = 0.5 g L⁻¹; HA = 5.0 mg L⁻¹; T = 25 °C.

3.5 Effect in simulated ground water

The decolorization of AO7 by activation of PS was studied in simulated ground water that contained a defined composition of inorganic and organic matter, including Fe(NO₃)₃·9H₂O (0.24 μM), NaHCO₃ (1.2 mM), Na₂SO₄ (0.34 mM), Na₂HPO4 (0.28 mM), NaCl (0.86 mM) and catechol (1 ppm).^8,44 As displayed in Fig. 8, the AO7 decolorization was significantly retarded in the simulated ground water compared to distilled water, especially in the first 60 min. After that, AO7 decolorized rapidly and reached 56.2% decolorization rate in 120 min. That was due to the scavenging reactions occurred between organic matters and SO₄⁻˙. On the other hand, the initial pH was changed to 6.83 when these inorganic and organic matters were added into distilled water to prepare simulated ground water, which declined the AO7 decolorization rate based on the above discussion.

Fig. 8 Decolorization of AO7 in simulated ground water. Experiment condition: PS = 2 mM; AO7 = 0.2 mM; ZVI = 0.5 g L⁻¹; T = 25 °C.

4. Conclusions

It was demonstrated that maximum AO7 decolorization occurred at pH 3.0. Increasing system pH resulted in a greater decrease in AO7 decolorization rates. AO7 decolorization efficiency was 100% at 120 min when the molar ratio of PS : AO7 was 5 : 1. The overall oxidation rate of AO7 was inhibited upon addition of NO₃⁻, NO₂⁻, SO₄²⁻, Cl⁻, CO₃²⁻, HCO₃⁻, HPO₄²⁻, H₂PO₄⁻ and EDTA, whereas ClO₄⁻, CH₃COO⁻ and HA were found to accelerate AO7 decolorization rates. ClO₄⁻, CH₃COO⁻ and HA with 50 mM, 10 mM and 1.0 mg L⁻¹, respectively were found to be the optimal concentration for AO7 decolorization. The other inorganic ions also exhibited different levels of impact on decolorizing of AO7, which was ranged as NO₂⁻ > H₂PO₄⁻ > HPO₄²⁻ > EDTA > SO₄²⁻ > CO₃²⁻ > HCO₃⁻ > NO₃⁻ > Cl⁻. The removal efficiencies of AO7 were decreased by 90.3%, 51.5% and 58.7%, respectively over 120 min in addition of NO₂⁻ (50 mM), H₂PO₄⁻ (50 mM) and HPO₄²⁻ (50 mM). While other inorganic ions including SO₄²⁻, CO₃²⁻, HCO₃⁻, NO₃⁻, Cl⁻ led to a decrease change of less than 20%.

On the basis of the above results and discussion, the reasons for the influence were as follows. (1) The ferric and ferrous ions underwent a complex reactions with H₂PO₄⁻, HPO₄²⁻, SO₄²⁻, CO₃²⁻, and HCO₃⁻, causing ions precipitation (such as H₂PO₄⁻ and HPO₄²⁻) to lose the active iron to activate PS and affecting the distribution of iron species. (2) Scavenging reactions occurred between inorganic ions and SO₄⁻˙ and formation of less reactive inorganic radicals (such as HCO₃˙ and CO₃⁻˙), which led to less SO₄⁻˙ towards AO7. (3) Oxidation reactions involving inorganic ions (such as NO₂⁻) consumed SO₄⁻˙. The mechanism of the acceleration by CH₃COO⁻ and HA was probably through acting as an electron ‘shuttle’ and facilitating electron transfer from ZVI surface to PS, which led to more Fe²⁺ releasing to the PS/ZVI systems to some extent. Furthermore, Fe²⁺ was continued to release and activate PS effectively during the reaction time. These findings indicate that CH₃COO⁻ and HA played an important role in activated PS applications. It was speculated that the AO7 decolorization enhancement caused by ClO₄⁻ was presumably ascribed to the oxidizing of ZVI directly by ClO₄⁻, which produced more Fe²⁺. The other anions and organic matters have been shown to affect wastewater treatment processes involved SO₄⁻˙. Activation of PS for degradation of refractory contaminants is a promising strategy in AOPs. The findings of this study will help achieve a deeper understanding of the impact of common inorganic ions and organic matters on the PS-based AOPs, which boosts the development of SO₄⁻˙-based AOPs.

Acknowledgements

National Science Foundation of China (Grant No. 31570568), State Key Laboratory of Pulp and Paper Engineering (201535), Guangdong High Level Tatent Project (201339).

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Footnote

† Electronic supplementary information (ESI) available. See DOI: 10.1039/c5ra16094d

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