Oxidative Degradation of Dimethyl Phthalate (DMP) by the Fe(VI

This is an open access article published under an ACS AuthorChoice License, which permits copying and redistribution of the article or any adaptations for non-commercial purposes.

Article Cite This: ACS Omega 2019, 4, 9467−9472

http://pubs.acs.org/journal/acsodf

Oxidative Degradation of Dimethyl Phthalate (DMP) by the Fe(VI)/ H2O2 Process Jinjuan Xue,* Zhengrong Zhu, Yuqing Zong, Chengjuan Huang, and Mingxin Wang* School of Environmental and Safety Engineering, Changzhou University, Changzhou 213164, P. R. China

Downloaded by 193.56.74.126 at 16:58:32:263 on May 30, 2019 from https://pubs.acs.org/doi/10.1021/acsomega.9b01012.

S Supporting Information *

ABSTRACT: This study investigates the degradation of dimethyl phthalate (DMP) with hydrogen peroxide and ferrate (Fe(VI)) under various reaction conditions. The results showed that the optimum conditions for dimethyl phthalate removal from water were as follows: (a) pH 7.0 and (b) the original molar ratio of [Fe(VI)]/[H2O2]/[DMP] equal to 10:2:1. Under the optimum conditions, the degradation rate of DMP can reach 89.7% in 360 min. Furthermore, 2,5dihydroxybenzaldehyde, isophthalic acid, 2-ethylhexanol, oxalic acid, 2,6-dihydroxybenzoic acid, 2,6-dihydroxybenzaldehyde, 2,5-dihydroxybenzoic acid, and monomethyl phthalate were identified as the degradation intermediates, and degradation pathways were proposed.

oxidation capacity, flocculation, disinfection, and deodorization function. The standard reduction potential in acidic pH is 2.20 V, whereas in basic pH this potential is about 0.72 V, and it is potentially the strongest of all of the oxidants/disinfectants realistically applicable to water and wastewater treatment.11−13 The reaction mechanism of Fe(VI) oxidation degradation of organic pollutants is electrophilic oxidation reaction including 1-e− and 2-e− processes, which have been reported in previous studies.14 Besides, Kamachi et al.15 proposed the increasing oxidizing power of the three types of Fe(VI) as FeO42− < HFeO4− < H2FeO4. Ohta et al.16,17 presented an argument through chemical calculations that there are two reaction mechanisms for the oxidation of methanol by Fe(VI), that is, addition−elimination and direct extraction. Fe(VI) can be converted into nanoscale ferric hydroxide during oxidation processes, which can be used as an alternative iron source for the Fenton reaction.18,19 The addition of hydrogen peroxide (H2O2) can reduce the amount of Fe(VI) and form Fentonlike reagents to further enhance the removal of target contaminants. The Fenton oxidation process is based on the nonselective oxidation ability of hydroxyl free radicals (HO•) generated from Fenton’s reagents. Owing to hydroxyl radicals’ high oxidation potential (E0 = 2.80 V), they can oxidize almost all of the organic substances and mineralize them to carbon dioxide and water. Its main reaction mechanism can be described as follows20,21

1. INTRODUCTION Phthalate esters (PAEs), also known as phthalates, are commonly used as plasticizers and in the manufacture of insecticide carriers, propellants, and cosmetics.1 Since PAEs are not chemically bound to the host plastics, they are inevitably leached from the products and released into the environment ultimately, and these organic compounds have caused a huge damage to water and air environment as pollutants. Dimethyl phthalate (DMP) ester is a common short-chained ester, which has been detected in many natural environments including marine waters, water courses, and sediments.2 DMP is relatively recalcitrant to biodegradation in the natural environment because of its benzene carboxylic group and its hydrolysis half-life of about 3 years.3 It has already been listed as a priority organic pollutant by the U.S. Environmental Protection Agency because of its toxicological properties and potential risks to human health, wildlife, and the environment.4 Thus, it is necessary to identify appropriate treatment technologies to remove DMP and remediate its adverse effects. Recently, several advanced technologies have been described for removing DMP, such as ozonation,5 photocatalysis,6,7 adsorption,8 chemical oxidation,9 and bioconversion.10 These technologies were shown to be capable of degrading various aqueous pollutants into less toxic substances. Among these possible treatment technologies, chemical oxidation has emerged as a promising method with a wide application due to its high efficiency, short reaction time, and simple operation. In recent years, the commonly used chemical oxidants have been Fenton’s reagent, ozone, potassium permanganate, potassium persulfate, and potassium ferrate. Compared with other chemical oxidants, Fe(VI) has been found to be an environment-friendly powerful oxidant over a wide pH range and a multifunctional water treatment agent because of its high © 2019 American Chemical Society

Fe2 + + H 2O2 → HO• + OH− + Fe3 +

(1)

Received: April 9, 2019 Accepted: May 16, 2019 Published: May 30, 2019 9467

DOI: 10.1021/acsomega.9b01012 ACS Omega 2019, 4, 9467−9472

ACS Omega HO• + organics → products + CO2 + H 2O

Article

(2)

However, using the traditional Fenton system has several drawbacks, such as easy precipitation of iron ions, loss of catalytic activity, and harsh environmental pH conditions.22,23 The Fe(VI)/H2O2 system can exert multiple functions, such as strong oxidation, formation of Fenton-like reagents, and flocculation and precipitation, to achieve efficient degradation of pollutants. The main interest of this study was to investigate the factors that affect the degradation efficiency of DMP by Fe(VI)/H2O2 and to identify the flocculation and degradation intermediates of DMP by the reaction system. The results show that Fe(VI)/ H2O2 advanced oxidation process is an efficient catalytic combustion technology for the destruction of DMP in water treatment.

Figure 2. Effect of the initial Fe(VI) concentration on DMP removal (pH = 7.0; molar ratio of [H2O2]/[DMP] = 2:1).

2. RESULTS AND DISCUSSION 2.1. Removal of DMP by a Single Fe(VI). To study the direct degradation of DMP by a single Fe(VI), five groups of experiments with different initial Fe(VI) concentrations were carried out and the results are shown in Figure 1. As can be

of Ct/Co in the presence of Fe(VI) at different concentrations as a function of the reaction time. The degradation of DMP generally can be regarded as a pseudo-first-order kinetics reaction: ln C t /C o = −k app t, where C o is the initial concentration of the DMP solution, Ct is the concentration of DMP at time t, and the slope kapp is the apparent reaction rate constant. From the experimental results shown in Figure 2, the DMP removal efficiency increases from 6.3 to 89.7% when increasing the molar ratio of [Fe(VI)]/[DMP] from 1:1 to 10:1, and the corresponding pseudo-first-order kinetics constant is improved from 0.00012 to 0.00623 min−1. However, Fe(VI) shows a slightly increased DMP removal efficiency of 4.2% when the molar ratio of Fe(VI)/[DMP] is up to 20:1. This phenomenon can be attributed to the shielding of OH• by excessive iron ions,30,31 as shown in eqs 3−5. Besides, the reduction of ferrate in water is caused by oxygen radicals and peroxide radicals. These radicals make ferrate become a strong oxidant and ferric ions make it a good coagulant. During oxidation, ferrate is reduced to Fe3+ or Fe(OH)3 and then flocculated. However, excessive Fe(VI) may lead to enhanced flocculation and coprecipitation of most organic high-molecular compounds.32

Figure 1. Degradation of DMP by a single Fe(VI).

HO• + Fe 2 + → OH− + Fe3 +

seen, the removal efficiency of DMP increases from 7.1 to 74.3% within the first 60 min when increasing the molar ratio of [Fe(VI)]/[DMP] from 2:1 to 20:1. A gradual further removal of DMP occurred in the next 120 min. The results show that DMP can be effectively removed by the single Fe(VI) system. Preliminary experimental results suggested that the concentration of DMP could not be effectively reduced by the single H2O2 system. H2O2 with a low redox potential was too easily decomposed to destroy the molecular structure of organic pollutants. However, DMP can be removed effectively by the single Fe(VI) system. Fe(VI) synergistically affected oxidization, coagulation, sterilization, and deodorization during the water treatment process, and its degradation product is iron hydroxide, which can further absorb oxidation byproducts. In addition, various studies indicate that Fe(VI) is a selective oxidizing agent that reacts with target compounds, such as phenol, chlorophenol, 3-methylphenol benzene series, methyl tert-butyl ether, acetaminophen, trichloroethylene, methanol, sulfite, and cyanide.24−29 In addition, the final product of Fe(VI) decomposition is ferric iron (Fe(III) solid), which may play an important role in the subsequent water treatment process, such as adsorption and coagulation. 2.2. Effect of the Initial Fe(VI) Concentration on the Degradation of DMP. Figure 2 shows the logarithmic graphs

(3)

HOO• + Fe 2 + → OOH + Fe3 + •

HOO + Fe

3+

→ Fe

2+

+ O2 + H

(4) +

(5)

2.3. Effect of the Initial H2O2 Concentration on the Degradation of DMP. The concentration of H2O2 is an important parameter in the advanced oxidation process of Fe(VI)/H2O2. The effect of the initial concentration of hydrogen peroxide on the degradation of DMP was studied with the molar ratio of [H2O2]/[DMP] varying from 0.25:1 to 5:1. As shown in Figure 3, the kapp increases with the increase of the concentration of hydrogen peroxide under the experimental conditions. The pseudo-first-order rate constant can reach up to 0.00623 min−1 when the molar ratio of [H2O2]/[DMP] is at levels of up to 2:1. However, Fe(VI)/ H2O2 exhibits a slight increase in the dimethylformamide removal efficiency when the molar ratio of [H2O2]/[DMP] rises from 2:1 to 5:1. In water, Fe(VI) is reduced to ferric iron (Fe(III)), which might serve as an in situ coagulant. The selfdecay of Fe(VI) to Fe(IV), Fe(V), and H2O2 was catalyzed by the Fe(III) precipitate,33 and H2O2 would further react with Fe(VI) to generate Fe(IV), which was also catalyzed by the Fe(III) precipitate.34 Lee et al.35 studied the decomposition of 9468


ACS Omega

Article

Fe(VI). The hydrogen peroxide anion (HO2−) is an ionic form of hydrogen peroxide, which is more reactive than hydrogen peroxide, and the alkaline condition is beneficial for the enhanced formation of hydroxyl radicals (eq 8).39 H 2O2 F HO2− + H+

(8)

HO2−

•

Meanwhile, is also a scavenger of HO (eq 9) and promotes the decomposition of H2O2 (eq 10).40 Furthermore, the self-decomposition rate of H2O2 also increases significantly with the increase of the solution pH.41,42 HO2− + HO• → HO2• + OH− −

HO2 + H 2O2 → H 2O + O2 + OH Figure 3. Effect of the initial H2O2 concentration on DMP removal (pH = 7.0; molar ratio of [Fe(VI)]/[DMP] = 10:1).

(9) −

(10)

4K 2FeO4 + 10H 2O → 4Fe(OH)3 + 8KOH + 3O2 ↑ (11)

Fe(VI) in phosphate-buffered water at pH 7.0 and determined that H2O2, Fe(III), and O2 were the main end products and that pseudo-first-order rate constants (reaction of ferrate(VI) with H2O2) increased linearly with increasing H2O2 concentration in the pH range of 7−12. These results indicate that the degradation of DMP in the experimental system may be controlled by the amount of hydroxyl radicals available. However, excessive hydrogen peroxide could be used as a scavenger of hydroxyl radicals, thus reducing the oxidation rate (eqs 6,7)36−38 HO• + HO• → H 2O2

(6)

H 2O2 + HO• → HO2• + H 2O

(7)

In addition, Fe(VI) is extremely unstable under strong acidic conditions (eq 11) and exhibits the greatest stability at about pH = 10.43 Rapid hydrolysis of Fe(VI) under strong acidic conditions and the extremely high stability of Fe(VI) under alkaline conditions will lead to low utilization of hydrogen peroxide, which is not conducive to the degradation of pollutants.44 These competing processes together control the concentration of hydroxyl radicals in the process of Fe(VI)/ H2O2 advanced oxidation and make kapp increase with the increase of the solution pH and then decrease gradually after reaching the peak value. 2.5. Effect of Coexisting Ions on the Degradation of DMP. Real wastewater commonly contains metal ions and inorganic anions; a solution containing the corresponding interfering ions (10 mg·L−1) was prepared and operated according to the recommended procedure, and the results are shown in Figure 5. From the experimental results, it can be

2.4. Effect of the Initial Solution pH on the Degradation of DMP. The solution pH plays an important role in the oxidation of organic contaminants by Fe(VI)/H2O2. The effect of the solution pH on the degradation of DMP was studied in the pH range of 3−11. As illustrated in Figure 4, the

Figure 5. Effect of initial coexisting ions on DMP removal (pH = 7.0; molar ratio of [Fe(VI)]/[H2O2]/[DMP] = 10:2:1). Figure 4. Effect of the initial solution pH on DMP degradation, where the molar ratio of [Fe(VI)]/[H2O2]/[DMP] =10:2:1.

seen that the effects of the six ions on DMP degradation vary. The DMP removal efficiency decreased slightly at different ion concentrations of K+ and SO42−. However, NO3−, Ca2+, HCO3−, and Mg2+ decreased the DMP removal efficiency significantly. The DMP removal efficiency decreased about 14.3, 35.1, 81.1, and 84.5% at 10 mg·L−1 of NO3−, Ca2+, HCO3−, and Mg2+, respectively. This can be attributed to the fact of that these ions can react with the hydroxyl radical easily, which may lead to a decrease of the DMP removal efficiency.45 In detail, bicarbonate may serve as a type of scavenger that can quench hydroxyl radicals (eq 12). Nitrate ions and sulfate ions can also capture and react with hydroxyl radicals, as shown in

oxidation efficiency of DMP is affected by the pH value because the acidity of the solution has a great influence on the Fenton reaction. Obviously, the kapp value increases with the increase of the pH value in the range of pH 3−5 and remains relatively stable in the pH range of 5−7. However, the reaction rate constant decreases significantly with the increase of the initial pH value under alkaline conditions. The effect of the solution pH on the efficiency of DMP degradation in the Fe(VI)/H2O2 system may be caused by three factors: hydroxyl radical production, hydroxyl radical scavenging, and stability of 9469


ACS Omega

Article

eq 13. In addition, the inhibiting effects of Mg2+ and Ca2+ were consistent with previous studies, indicating that hydroxide radicals would adhere to the surface of ferrite and have adverse effects on iron oxidation.46 SO4 2 − + HO• → SO4•− + OH−

(12)

HCO3− + HO• → CO3•− + H 2O

(13)

4. EXPERIMENTAL SECTION 4.1. Chemicals and Reagents. Analytical-grade DMP, hydrogen peroxide (30% w/w), magnesium chloride, sodium hydroxide, sodium bicarbonate, potassium chloride, sodium nitrate, sodium sulfite, calcium chloride, sulfuric acid, sodium sulfate, hydrochloric acid, acetic acid, sodium acetate, hydroxylamine hydrochloride, and ferrous ammonium sulfate were purchased from Sinopharm Chemical Reagents (Shanghai, China). Solid potassium ferrate (K2FeVIO4) and potassium bromide (spectrography) were purchased from MACKLIN Reagent (Shanghai, China). Methanol and acetonitrile (HPLC grade) were purchased from Merck (Darmstadt, Germany). Methyl tert-butyl ether (GC grade), the GC derivatization reagent, hexamethyldisilazane, and trimethylchlorosilane were purchased from Aladdin Reagent (Shanghai, China). Ultrapure water obtained from a Milli-Q water purification system was used in the experiment. The DMP stock solution (100 mg·L−1) was prepared in ultrapure water. 4.2. DMP Degradation Experiments. Kinetic experiments of DMP degradation with Fe(VI)/H 2 O 2 were conducted in 250 mL Erlenmeyer flasks with stoppers in a thermostatic shaker. Varying doses of Fe(VI), H2O2, and coexisting ions were added in 100 mL of DMP aqueous solution (100 mg·L−1). After the reaction started, 1.0 mL of the DMP sample was collected into a chromatography vial containing 200 μL of sodium sulfite (0.10 M) to quench the free radical at reaction time intervals of 30, 60, 90, 120, 180, 240, 300, and 360 min for the HPLC analysis. The pH of test solutions was adjusted with 0.01 M H2SO4 and 0.01 M NaOH. All samples in the experiments were filtered using 0.45 μm membrane filters. Each test was conducted twice in the course of the experiment. The average values of two DMP concentrations were taken as the current DMP concentrations to calculate the reaction rate in each test. 4.3. Identification of Degradation Products. Then, 1.0 mL of the DMP sample was collected after 360 min and concentrated by vacuum freeze-drying. The residue was dissolved in 4 mL of methyl tertiary butyl ether and then derivatized by 0.2 mL of hexamethyldisilazane and 0.1 mL of trimethylchlorosilane at 70 °C for 30 min before it was filtered using 0.45 μm membrane filters. GC/MS was utilized to analyze the degradation products. 4.4. Analytical Methods. The concentration of DMP was determined using Shimadzu LC-20AT HPLC equipped with a UV detector. The column utilized to separate the analytes was a C18 reversed-phase column SinChrom ODS-BP (5 μm, 4.6 × 250 mm2). The mobile phase consisted of water (30%) and methanol (70%) and was eluted at 1.0 mL·min−1. The injection volume was 20 mL and the detection wavelength was 225 nm. The degradation products were detected using Thermo Trace1300 GC−MS with a TG-5 MS capillary column. Nitrogen was the carrier gas with a flow rate of 1 mL·min−1. The injector port was held at 280 °C. The following procedure was applied to the column: an initial oven temperature of 80 °C for 3 min, a ramp to 120 °C at 30 °C· min−1, then increasing to 140 °C at 18 °C·min−1, followed by a ramp to 250 °C at 30 °C·min−1, and finally maintained at 250 °C for 1 min.

2.6. Degradation Intermediates and Degradation Mechanism. Based on the identified intermediates in Figure S1 and Table S1, the proposed degradation pathways are shown in Figure 6. Methyl phthalate was formed through

Figure 6. Potential DMP degradation pathways.

direct oxidation and hydrolysis of esters on the side of DMP. In a similar way, another methoxy group was lost to form phthalic acid, which is further converted into isophthalic acid by isomerization. Frontier electron density analysis shows that the benzene ring is more vulnerable to free radical attack when the number of carbon atoms in the alkyl chain is less than three.47 Thus, benzoic acid was attacked by free radicals to produce 2,5-dihydroxybenzoic acid and 2,6-dihydroxybenzoic acid, which had been confirmed by a prior report.48 Subsequently, 2,5-dihydroxybenzaldehyde and 2,6-dihydroxybenzaldehyde are formed by further oxidation. The generated intermediates were further decomposed to oxalic acid or other small molecular organic acids and alcohols. Ultimately, the mineralization of DMP was completely accomplished until almost all of the formed intermediates decompose to CO2 and H2O.

3. CONCLUSIONS This study illustrates that Fe(VI)/H2O2 can effectively decompose and mineralize DMP in aqueous solutions. The DMP removal efficiency was influenced by the Fe(VI) concentration, H2O2 concentration, pH, and coexisting ions. The optimal cost-effective conditions for the degradation of DMP in the Fe(VI)/H2O2 system were observed at pH 7.0, with an initial molar ratio of [Fe(VI)]/[H2O2]/[DMP] = 10:2:1. Also, Mg2+ and HCO3− in aqueous solutions could significantly decrease the DMP degradation efficiency. Moreover, eight degradation intermediates were identified by gas chromatography−mass spectrometry (GC/MS), and possible degradation pathways were proposed. 9470


ACS Omega

■

Article

pyridines using novel pyrazines-degrading microorganisms. Chemija 2013, 24, 67−73. (11) Liu, H.; Chen, J.; Wu, N.; Xu, X.; Qi, Y.; Jiang, L.; Wang, X.; Wang, Z. Oxidative degradation of chlorpyrifos using ferrate(VI): Kinetics and reaction mechanism. Ecotoxicol. Environ. Saf. 2019, 170, 259−266. (12) Jiang, J.; Stanford, C.; Petri, M. Practical application of ferrate (VI) for water and wastewater treatment-Site study’s approach. Water Energy Nexus 2018, 1, 42−46. (13) Bonakdarpour, A.; Cataldo, M.; Mohseni, M.; Wilkinson, D. In Electrosynthesis of Near-Neutral Ferrate Species for Drinking Water Treatment Using a Recirculating Batch Reactor, Meeting Abstracts, The Electrochemical Society, 2018; p 1300. (14) Sharma, V. Oxidation of inorganic compounds by ferrate (VI) and ferrate (V): one-electron and two-electron transfer steps. Environ. Sci. Technol. 2010, 44, 5148−5152. (15) Kamachi, T.; Nakayama, T.; Yoshizawa, K. Mechanism and kinetics of cyanide decomposition by ferrate. Bull. Chem. Soc. Jpn. 2008, 81, 1212−1218. (16) Ohta, T.; Kamachi, T.; Shiota, Y.; Yoshizawa, K. A theoretical study of alcohol oxidation by ferrate. J. Org. Chem. 2001, 66, 4122− 4131. (17) Sharma, V. Ferrate (VI) and ferrate (V) oxidation of organic compounds: kinetics and mechanism. Coord. Chem. Rev. 2013, 257, 495−510. (18) Anquandah, G.; Sharma, V.; Knight, D.; Batchu, S.; Gardinali, P. Oxidation of trimethoprim by ferrate (VI): kinetics, products, and antibacterial activity. Environ. Sci. Technol. 2011, 45, 10575−10581. (19) Casbeer, E.; Sharma, V.; Zajickova, Z.; Dionysiou, D. Kinetics and mechanism of oxidation of tryptophan by ferrate (VI). Environ. Sci. Technol. 2013, 47, 4572−4580. (20) Wang, Y.; Liu, Y.; Liu, T.; Song, S.; Gui, X.; Liu, H.; Tsiakaras, P. Dimethyl phthalate degradation at novel and efficient electroFenton cathode. Appl. Catal., B 2014, 156−157, 1−7. (21) He, P.-J.; Zheng, Z.; Zhang, H.; Shao, L.; Tang, Q. PAEs and BPA removal in landfill leachate with Fenton process and its relationship with leachate DOM composition. Sci. Total Environ. 2009, 407, 4928−4933. (22) Zhao, H.; Qian, L.; Chen, Y.; Wang, Q.; Zhao, G. Selective catalytic two-electron O2 reduction for onsite efficient oxidation reaction in heterogeneous electro-Fenton process. Chem. Eng. J. 2018, 332, 486−498. (23) Sanchis, S.; Meschede-Anglada, L.; Serra, A.; Simon, F.; Sixto, G.; Casas, N.; Garcia-Montano, J. Solar photo-Fenton with simultaneous addition of ozone for the treatment of real industrial wastewaters. Water Sci. Technol. 2018, 77, 2497−2508. (24) Graham, N.; Jiang, C.; Li, X.; Jiang, J.; Ma, J. The influence of pH on the degradation of phenol and chlorophenols by potassium ferrate. Chemosphere 2004, 56, 949−956. (25) Wang, G.; Feng, L. Experimental studies on application of potassium ferrate for 3-methylphenol removal from solution in laboratory. Procedia Environ. Sci. 2013, 18, 486−492. (26) Minetti, R. C. P.; Macaño, H. R.; Britch, J.; Allende, M. C. In situ chemical oxidation of BTEX and MTBE by ferrate: pH dependence and stability. J. Hazard. Mater. 2017, 324, 448−456. (27) Wang, H.; Liu, Y.; Jiang, J. Reaction kinetics and oxidation product formation in the degradation of acetaminophen by ferrate (VI). Chemosphere 2016, 155, 583−590. (28) Dobosy, P.; Vizsolyi, É ; Varga, I.; Varga, J.; Láng, G.; Záray, G. Trichloroethylene removal from water by ferrate treatment. Microchem. J. 2016, 127, 74−78. (29) Gombos, E.; Barkács, K.; Felföldi, T.; Vértes, C.; Makó, M.; Palkó, G.; Záray, G. Removal of organic matters in wastewater treatment by ferrate (VI)-technology. Microchem. J. 2013, 107, 115− 120. (30) Jiang, J.; Zhou, Z.; Pahl, O. Preliminary study of ciprofloxacin (cip) removal by potassium ferrate(VI). Sep. Purif. Technol. 2012, 88, 95−98.

ASSOCIATED CONTENT

* Supporting Information S

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acsomega.9b01012. GC−MS chromatograms of the intermediates during the DMP degradation; intermediate products of DMP by Fe(VI) and Fe(VI)/H2O2 (PDF)

■

AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected] (J.X.). *E-mail: [email protected] (M.W.). ORCID

Jinjuan Xue: 0000-0001-8394-1372 Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS This work was supported by the National Natural Science Foundation of China (grant nos. 21808019, 41641032, and 41772240), the Natural Science Foundation of Jiangsu (grant no. BK20180958), the Science and Technology Bureau of Changzhou (CJ20179037), and Science and Technology Project of Changzhou University (ZMF17020042).

■

REFERENCES

(1) Liu, H.; Cui, K.; Zeng, F.; Chen, L.; Cheng, Y.; Li, H.; Li, S.; Zhou, X.; Zhu, F.; Ouyang, G.; Luan, T.; Zeng, Z. Occurrence and distribution of phthalate esters in riverine sediments from the Pearl River Delta region, South China. Mar. Pollut. Bull. 2014, 83, 358− 365. (2) Yuan, B.; Li, X.; Graham, N. Aqueous oxidation of dimethyl phthalate in a Fe(VI)-TiO2-UV reaction system. Water Res. 2008, 42, 1413−1420. (3) He, W.; Qin, N.; Kong, X.; Liu, W.; He, Q.; Ouyang, H.; Yang, C.; Jiang, Y.; Wang, Q.; Yang, B.; Xu, F. Spatio-temporal distributions and the ecological and health risks of phthalate esters (PAEs) in the surface water of a large, shallow Chinese lake. Sci. Total Environ. 2013, 461−462, 672−680. (4) Osman, B.; Ozer, E.; Demirbel, E.; Gucer, S.; Besirli, N. Synthesis and characterization of L-tryptophan containing microbeads for removal of dimethyl phthalate from aqueous phase. Sep. Purif. Technol. 2013, 109, 40−47. (5) Liu, Y.; Wu, D.; Peng, S.; Feng, Y.; Liu, Z. Enhanced mineralization of dimethyl phthalate by heterogeneous ozonation over nanostructured Cu-Fe-O surfaces: Synergistic effect and radical chain reactions. Sep. Purif. Technol. 2019, 209, 588−597. (6) Wang, G.; Chen, Q.; Liu, Y.; Ma, D.; Xin, Y.; Ma, X.; Zhang, X. In situ synthesis of graphene/WO3 co-decorated TiO2 nanotube array photoelectrodes with enhanced photocatalytic activity and degradation mechanism for dimethyl phthalate. Chem. Eng. J. 2018, 337, 322−332. (7) Ku, Y.; Shiu, S.; Wu, H. Decomposition of dimethyl phthalate in aqueous solution by UV-LED/TiO2 process under periodic illumination. J. Photochem. Photobiol., A 2017, 332, 299−305. (8) Zhao, H.; Wang, Q.; Chen, Y.; Tian, Q.; Zhao, G. Efficient removal of dimethyl phthalate with activated iron-doped carbon aerogel through an integrated adsorption and electro-Fenton oxidation process. Carbon 2017, 124, 111−122. (9) Wang, Z.; Deng, D.; Yang, L. Degradation of dimethyl phthalate in solutions and soil slurries by persulfate at ambient temperature. J. Hazard. Mater. 2014, 271, 202−209. (10) Kutanovas, S.; Rutkiene, R.; Urbelis, G.; Tauraite, D.; Stankeviciute, J.; Meskys, R. Bioconversion of methylpyrazines and 9471


ACS Omega

Article

(31) Wang, H.; Zhao, Y.; Su, Y.; Li, T.; Yao, M.; Qin, C. Fenton-like degradation of 2,4-dichlorophenol using calcium peroxide particles: performance and mechanisms. RSC Adv. 2017, 7, 4563−4571. (32) Yuan, B.; Qu, J.; Fu, M. Removal of cyanobacterial microcystinLR by ferrate oxidation−coagulation. Toxicon 2002, 40, 1129−1134. (33) Jiang, Y.; Goodwill, J.; Tobiason, J.; Reckhow, D. Effect of different solutes, natural organic matter, and particulate Fe (III) on ferrate (VI) decomposition in aqueous solutions. Environ. Sci. Technol. 2015, 49, 2841−2848. (34) Jiang, Y.; Goodwill, J.; Tobiason, J.; Reckhow, D. Bromide oxidation by ferrate (VI): The formation of active bromine and bromate. Water Res. 2016, 96, 188−197. (35) Lee, Y.; Kissner, R.; von Gunten, U. Reaction of ferrate (VI) with ABTS and self-decay of ferrate (VI): kinetics and mechanisms. Environ. Sci. Technol. 2014, 48, 5154−5162. (36) Tian, F.; Ma, S.; Xu, B.; Hu, X.; Xing, H.; Liu, J.; Wang, J.; Li, Y.; Wang, B.; Jiang, X. Photochemical degradation of iodate by UV/ H2O2 process: Kinetics, parameters and enhanced formation of iodotrihalomethanes during chloramination. Chemosphere 2019, 221, 292−300. (37) Guo, H.; Gao, N.; Chu, W.; Li, L.; Zhang, Y.; Gu, J.; Gu, Y. Photochemical degradation of ciprofloxacin in UV and UV/H2O2 process: kinetics, parameters, and products. Environ. Sci. Pollut. Res. 2013, 20, 3202−3213. (38) Xu, B.; Gao, N.; Cheng, H.; Xia, S.; Rui, M.; Zhao, D. Oxidative degradation of dimethyl phthalate (DMP) by UV/H2O2 process. J. Hazard. Mater. 2009, 162, 954−959. (39) Beltrán, F. J.; Ovejero, G.; Rivas, J. Oxidation of polynuclear aromatic hydrocarbons in water. 3. UV radiation combined with hydrogen peroxide. Ind. Eng. Chem. Res. 1996, 35, 883−890. (40) Suganuma, R.; Kawasaki, S.; Yasuoka, K. Disinfection by HO2 radicals generated by negative corona discharge on acidic solution. IEEE Trans. Plasma Sci. 2019, 47, 346−349. (41) Lin, C.; Lin, H.; Hsu, L. Degradation of ofloxacin using UV/ H2O2 process in a large photoreactor. Sep. Purif. Technol. 2016, 168, 57−61. (42) Zhou, C.; Gao, N.; Deng, Y.; Chu, W.; Rong, W.; Zhou, S. Factors affecting ultraviolet irradiation/hydrogen peroxide (UV/ H2O2) degradation of mixed N-nitrosamines in water. J. Hazard. Mater. 2012, 231−232, 43−48. (43) Daneshvar, N.; Behnajady, M.; Mohammadi, M.; Dorraji, M. UV/H2O2 treatment of Rhodamine B in aqueous solution: Influence of operational parameters and kinetic modeling. Desalination 2008, 230, 16−26. (44) Dong, H.; Qiang, Z.; Liu, S.; Li, J.; Yu, J.; Qu, J. Oxidation of iopamidol with ferrate (Fe(VI)): Kinetics and formation of toxic iodinated disinfection by-products. Water Res. 2018, 130, 200−207. (45) Kalsoom, U.; Ashraf, S.; Meetani, M.; Rauf, M.; Bhatti, H. Degradation and kinetics of H2O2 assisted photochemical oxidation of Remazol Turquoise Blue. Chem. Eng. J. 2012, 200−202, 373−379. (46) Ahmed, T.; Uddin, F.; Azmat, R. Kinetics and mechanism study of chemical treatment of methylene green by urea. Chin. J. Chem. 2010, 28, 748−754. (47) Yuan, B.; Li, X.; Graham, N. Reaction pathways of dimethyl phthalate degradation in TiO2-UV-O2 and TiO2-UV-Fe(VI) systems. Chemosphere 2008, 72, 197−204. (48) Liu, X.; Zhang, T.; Zhou, Y.; Fang, L.; Shao, Y. Degradation of atenolol by UV/peroxymonosulfate: kinetics, effect of operational parameters and mechanism. Chemosphere 2013, 93, 2717−2724.

9472


Oxidative Degradation of Dimethyl Phthalate (DMP) by the Fe(VI

Recommend Documents