Article pubs.acs.org/est
Selective Oxidation of Arsenite by Peroxymonosulfate with High Utilization Efficiency of Oxidant Zhaohui Wang,*,†,‡ Richard T. Bush,† Leigh A. Sullivan,† Chuncheng Chen,*,§ and Jianshe Liu‡ †
Southern Cross GeoScience, Southern Cross University, Lismore, NSW 2480, Australia College of Environmental Science and Engineering, Donghua University, Shanghai, 201620, China § Beijing National Laboratory for Molecular Sciences, Key Laboratory of Photochemistry, Institute of Chemistry, The Chinese Academy of Sciences, Beijing 100190, China ‡
S Supporting Information *
ABSTRACT: Oxidation of arsenite (As(III)) is a critical yet often weak link in many current technologies for remediating contaminated groundwater. We report a novel, efficient oxidation reaction for As(III) conversion to As(V) using commercial available peroxymonosulfate (PMS). As(III) is rapidly oxidized by PMS with a utilization efficiency larger than 90%. Increasing PMS concentrations and pH accelerate oxidation of As(III), independent to the availability of dissolved oxygen. The addition of PMS enables As(III) to oxidize completely to As(V) within 24 h, even in the presence of high concentrations of radical scavengers. On the basis of these observations and theoretical calculations, a two-electron transfer (i.e., oxygen atom transfer) reaction pathway is proposed. Direct oxidation of As(III) by PMS avoids the formation of nonselective reactive radicals, thus minimizing the adverse impact of coexisting organic matter and maximizing the utilization efficiency of PMS. Therefore, this simple approach is considered a cost-effective water treatment method for the oxidation of As(III) to As(V).
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INTRODUCTION Arsenic is an element that raises much concern over environmental quality and human health. The greatest threat to public health from arsenic originates from contaminated groundwater.1 As(III) and As(V) are the two most common naturally occurring oxidation states of dissolved arsenic, with As(III) exceedingly more harmful due to its much higher mobility and toxicity.2 Therefore, oxidizing As(III) to As(V) is potentially an effective strategy to reduce the impacts of arsenic on society. In addition, its oxidation is usually a prerequisite step for most subsequent arsenic removal technologies like coagulation, sorption, and membrane filtration.3 Overcoming the slow oxidation rate of As(III) in airsaturated solutions is one of the major technical challenges.4 Numerous methods based on the advanced oxidation processes (AOPs) including Fenton (like) reaction,5 zerovalent iron oxidation6,7 and photocatalysis,8 have been developed to enhance the kinetics of As(III) oxidation. While these advanced AOPs are effective for the oxidation of As(III), there are significant inherent limitations to their practical application for the remediation of contaminated groundwater. For example, all these oxidation reactions are known to be initiated by the nonselective radicals such as •OH radicals.9 These reactive radicals tend to oxidize the common coexisting nontoxic organic matter species such as humic acids, besides the targeted As(III), which leads to much lower utilization efficiency (UE) of the oxidant. In addition, the use of UV irradiation or addition © 2014 American Chemical Society
of excessive amounts of metal catalysts (e.g., Fe(II)) is evidently unfavorable for the treatment of As(III)-contaminated groundwater.5,8 The oxidation of pollutants by using peroxydisulfate (S2O82−) and peroxymonosulfate (HSO5−, PMS) as the main parent oxidants is one of the emerging AOPs that is gaining importance in water treatment applications.10−13 In these systems, the dominant active species was conventionally considered to be the sulfate radicals (SO4•−, E0 = 2.5−3.1 V vs NHE), which are formed through activation of the peroxosulfates by heat,14 UV radiation,15,16 or transition metal catalysts10,17 (eq 1). S2 O82 − /HSO5− + initiator → SO4•− + (HSO5• , SO4 2 − ,• OH)
(1)
Peroxosulfate-based AOPs have recently been introduced into the oxidation of As(III). The UV/peroxydisulfate and Fe(II)/peroxydisulfate systems, for example, can rapidly oxidize the As(III) over a broad pH range (2.0−8.0).16 In these systems, the oxidation of As(III) is apparently initiated by highly reactive sulfate radicals with As(IV) as the intermediate Received: Revised: Accepted: Published: 3978
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improved. In addition, the direct oxidation without the use of UV light, other transition metals, or heating is favorable for its application for the practical remediation of As-contaminated groundwater.
(eq 2), and the As(IV) further reacts with molecular oxygen (eq 3) or another sulfate radical to generate As(V) (eq 4). SO4•− + As(III) → As(IV) + SO4 2 −
(2)
O2 + As(IV) → As(V) + O4•−
(3)
SO4•− + As(IV) → As(V) + SO4 2 −
(4)
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EXPERIMENTAL SECTION Chemicals. Oxone ([2KHSO5·KHSO4·K2SO4] salt, 95%) and ammonium molybdate tetrahydrate ((NH4)6Mo7O24· 4H2O, 99%) were purchased from Sigma-Aldrich. Sodium arsenite (NaAsO2 ≥ 97%) was purchased from Ajax Finechem. Iron(II) sulfate (FeSO4·7H2O, >99%) was obtained from Acros Organics. Pahokee peat humic acid standard was purchased from International Humic Substances Society. Sodium arsenate (Na2HAsO4·7H2O), NaCl, Na2SO4, NaF, KSCN, NaOH, H2SO4, HCl, methanol, 2-propanol, ascorbic acid, and antimony potassium tartrate (K(SbO)C4H4O6·0.5H2O) were of reagent grade and used as supplied. Milli-Q UltraPure water (18.3 MΩ cm) was used for all experiments. Experimental Procedures. A common stock solution of each reactant was prepared. Aliquots of the stock solutions were combined to achieve the initial experimental conditions. The experiments commenced by the addition of PMS. Unless otherwise noted, reactions were carried out at room temperature (20 ± 2 °C) under exposure to air. The reaction vessel was mantled by an inverted 500 mL beaker wrapped with aluminum foil to minimize the photolysis of PMS by ambient light. The pH’s of reaction solutions were not controlled (pHi = 4.1) except for the pH effect experiments. The pH drift was always NaCl > NaSO4 (Figure 4b). As the cations are the same in the three salts, the different effect should be attributed to the anions. It is interesting that PMS can directly oxidize chloride to form active chlorine species31 (i.e., Cl2/ClO−) which can further oxidize arsenite.32 In addition, the effect of salts appeared to increase with their concentrations (Figures S7 and S8), suggesting that the
Figure 4. Effects of (a) pH and (b) inorganic anions on the oxidation of As(III). Conditions a: As(III), 20 μM; PMS, 50 μM. Conditions b: As(III), 40 μM; PMS, 200 μM; pH 4.1.
inorganic anions are involved in PMS activation and thereby in As(III) oxidation by PMS. Mechanistic Discussion. One- versus Two-Electron Oxidation Process. Our results indicate two distinguished processes for the oxidation of As(III) by PMS: rapid process I and mild process II. The present observations, such as less selectivity for As(III), lower UEs of PMS, the enhancement of oxidation by Fe(II) and O2, and quenching by radical scavengers, indicate that process I, which can account for 30% of overall oxidation, is a typical free-radical-based reaction initiated by the single-electron reduction of PMS. During the oxidation of As(III) by PMS, 70% of the reaction is via process II, which proceeds mildly, with high selectivity and UEs of PMS. This process is independent of the presence of a scavenger, indicative of the nonradical characteristics of this process. Process II may proceed through a direct oxidation (oxygen atom transfer) mechanism.33 This pathway involves heterolytic cleavage of the peroxo bond and an oxygen atom transfer, converting PMS itself to the sulfate ion upon reaction, as shown in eq 12, where A is a reagent with an oxygen acceptor site such as As(III). HSO−5 + A → HSO−4 + OA
(12)
One- versus two-electron oxidation pathways can be predicted by the use of reduction potentials of reactants and products. The standard potential of HSO5−/HSO4− is 1.85 V vs NHE, whereas that of SO4•−/HSO4− is 2.49 V vs NHE.19,26 The potential for one-electron reduction of HSO5− can be 3982
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calculated to be 1.21 V vs NHE for the redox couple HSO5−/ SO4•−. The one-electron oxidation of As(III) would involve the formation of the As(IV) intermediate, and the corresponding redox potential of AsIV(OH)4/AsIII(OH)3 (2.4 V vs NHE) is much higher than 1.21 V (Table S2). Therefore, the formation of As(IV) species and SO4•− might be a thermodynamically uphill step in this one-electron pathway. In the case of Fe(II) (E0(Fe3+/Fe2+) = 0.77 V vs NHE),20 a one-electron first step to generate sulfate radicals is thermodynamically favorable. Anipsitakis and Dionysiou reported that Ce(IV) and Fe(III) can both decompose PMS to HSO5•.17 Since As(III) cannot act as an electron acceptor like Ce(IV) and Fe(III), we expect that the HSO5• will not form via the reaction of PMS with As(III). The redox potential of two electron oxidation of As(III) (E0(H3AsVO4/H3AsIIIO3)) is 0.56 V vs NHE (Table S2), implying As(III) is readily oxidized by PMS (E0(HSO5−/ HSO4−) = 1.85 V vs NHE) via a two-electron transfer pathway. The two-electron oxidation process of As(III) by PMS is also supported by the present experimental data: (1) The rate of As(III) oxidation was less impeded by methanol and 2propanol in As(III)/PMS systems, as compared to Fe(II)/ As(III)/PMS systems, indicating that As(III) participated differently than the one-electron pathway of PMS activation by Fe(II).17 (2) Methylene blue dye, a commonly used indicator for reactive radicals formation (e.g., •OH, SO4•−), was not decolored in the As(III)/PMS system, while a further addition of Fe(II) led to dye decoloration (Figure 3). This provides additional evidence for the nonradical characteristics of the overall As(III)/PMS reaction. The proposed two-electron transfer (oxygen atom transfer) mechanism was further verified by DFT calculation (see calculation details in the Supporting Information). Under the present experimental conditions (pH = 4.1), PMS should be in monoanionic form (HSO5−), while As(III) should be present neutrally (As(OH)3).2 Prior to the oxygen transfer, a reactant complex is formed through two hydrogen bonds between two protons of the As(OH)3 and two oxygen atoms from HSO5−, with an energy decrease of 7.0 kcal/mol relative to isolated reactants. The formation of double hydrogen bonds suggests that the two reactants As(OH)3 and HSO5− are facile to interact with each other, which will preferentially favor the selective oxidation of As(III) by the PMS. In the transition state, the calculation shows the transferred oxygen atom moves with the proton together, along with cleavage of the O−O bond of PMS. Activation energy relative to the reactant complex is 18.0 kcal/mol (Figure 5), indicating that this reaction can proceed smoothly at room temperature. In the product complex, the terminal oxygen of the PMS is transferred to the vacant site of As(III), and the oxidized AsO(OH)3 is formed. In the absence of any proton acceptor, the peroxide proton is left to the residual oxygen after the cleavage of the O−O bond of PMS to form HSO4−. The overall reaction energy is 50.6 kcal/mol, which is consistent with the above thermodynamic estimation that the two-electron oxidation of As(III) by PMS is a strong exothermic reaction. As the oxidation reaction was carried out in solution, the H2O molecule was explicitly included in the reaction as a hydrogen acceptor. It was observed that the presence of a H2O molecule slightly increases the activation energy (18.9 kcal/ mol) and decreases the reaction energy (47.2 kcal/mol), because of the weak proton acceptor ability of H2O. The effect of the pH and inorganic anions on the oxidation rate of As(III) (Figure 4) also prompts us to include the simple anions Cl−,
Figure 5. Energy profiles calculated by DFT (B3LYP functional and mixed basis sets cc-pVTZ-PP for As and cc-pVTZ for all other atoms) for the oxygen transfer from PMS to As(OH)3 with respect to the reactant complexes (RC) in the presence of various proton acceptors. TS represents transition states, and PC is the product complex. The upper illustrates the obtained PC, TS, and PC structures for the reaction with a typical proton acceptor (F−), in which the oxygen and proton of the peroxide are transferred to the As(III) and proton acceptor, respectively.
F−, and OH− in the calculation. The presence of anions as hydrogen acceptors Cl−, F−, and OH− is found to largely decrease the activation energy of the oxygen atom transfer by interacting with the proton of the peroxide, in agreement with the enhanced oxidation rate in the presence of inorganic salts. The effect is largely dependent on the ability for the anion to accept a proton and in the order of OH− > F− > Cl−. In the presence of OH− and F− (strong proton acceptors), the activation energies decrease from 18.9 kcal/mol with a H2O molecule as an acceptor to 5.3 and 9.7 kcal/mol, respectively, which is corresponding to the experimental observation of enhanced oxidation of As(III) at high pH and after adding F−. The self-decomposition of PMS at alkaline pH may produce oxygen and sulfate but does not necessarily contribute to the oxidation of As(III).34 As a weak proton acceptor, Cl− only slightly reduces the activation energy (17.6 kcal/mol), suggesting that a high concentration of Cl− would be required to achieve the same enhancing effect as the OH− and F− (Figure S7). Further, the exothermicity of oxidation is also enhanced along with activation energy by the same magnitude. These theoretical results are in agreement with the experimental observations on the positive effects of increasing pH and inorganic anions on kinetics of As oxidation (Figure 5), which verifies the feasibility of the proposed oxygen transfer (two-electron oxidation) pathway. Technical Implications. The determination of As(III) oxidation rates by varying concentrations of PMS and As(III), pH, and inorganic anions allows general predictions about the applicability of PMS oxidant for the oxidative transformation of As(III) in groundwater. We provide primary evidence that indicates the PMS/As(III) reaction is an attractive new alternative remediation technology for municipal or central water treatment. It operates across a broad range of geochemical conditions, without needing the addition of transition-metal catalysts,17 energy inputs (e.g., heat, light, and ultrasonic),14−16,35 or pH adjustment. With respect to the oxidant, PMS is cost-efficient ([PMS]/As(III) = 1:1) and easy to handle (solid form), and its decomposition to sulfate is 3983
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normally considered nonpolluting.15 Furthermore, the PMS/ As(III) reaction is not unduly affected by the presence of degradable organic compounds. This contrasts with many of the other As(III) oxidation technologies for water treatment, such as Fenton reaction5,23 and zerovalent iron activation.28 The efficiency of the coupled PMS/As(III) reaction seems to depend on several other factors such as pH, inorganic anions, and coexisting metal ions as well. Therefore, a thorough understanding of water quality parameters in As(III)-contaminated groundwater is a prerequisite for the field application of PMS in order to predict the As(III) oxidation efficiency.
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(10) Anipsitakis, G. P.; Dionysiou, D. D. Degradation of organic contaminants in water with sulfate radicals generated by the conjunction of peroxymonosulfate with cobalt. Environ. Sci. Technol. 2003, 37, 4790−4797. (11) Xu, L.; Yuan, R.; Guo, Y.; Xiao, D.; Cao, Y.; Wang, Z.; Liu, J. Sulfate radical-induced degradation of 2,4,6-trichlorophenol: a de novo formation of chlorinated compounds. Chem. Eng. J. 2013, 217, 169− 173. (12) Chan, K. H.; Chu, W. Degradation of atrazine by cobaltmediated activation of peroxymonosulfate: different cobalt counteranions in homogenous process and cobalt oxide catalysts in photolytic heterogeneous process. Water Res. 2009, 43, 2513−2521. (13) Yuan, R.; Ramjaun, N. S.; Wang, Z.; Liu, J. Effects of chloride ion on degradation of Acid Orange 7 by sulfate radical-based advanced oxidation process: Implications for formation of chlorinated aromatic compounds. J. Hazard. Mater. 2011, 196, 173−179. (14) Liang, C. J.; Su, H.-W. Identification of sulfate and hydroxyl radicals in thermally activated persulfate. Ind. Eng. Chem. Res. 2009, 48, 5558−5562. (15) Lau, T. K.; Chu, W.; Graham, N. J. D. The aqueous degradation of butylated hydroxyanisole by UV/S2O82−: Study of reaction mechanisms via dimerization and mineralization. Environ. Sci. Technol. 2007, 41, 613−619. (16) Neppolian, B.; Celik, E.; Choi, H. Photochemical oxidation of arsenic(III) to arsenic(V) using peroxydisulfate ions as an oxidizing agent. Environ. Sci. Technol. 2008, 42, 6179−6184. (17) Anipsitakis, G. P.; Dionysiou, D. D. Radical generation by the interaction of transition metals with common oxidants. Environ. Sci. Technol. 2004, 38, 3705−3712. (18) Liang, C. J.; Lee, I.-L.; Hsu, I.-Y.; Liang, C.-P.; Lin, Y.-L. Persulfate oxidation of trichloroethylene with and without iron activation in porous media. Chemosphere 2008, 70, 426−435. (19) DuPont Oxone® Monopersulfate Compound. http://www2. dupont.com/Oxone/en_US/assets/downloads/K20102_Oxone_ Technical_Bulletin.pdf. (accessed Sept. 21, 2013). (20) Bard, A. J.; Parsons, R.; Jordan, J. Standard Potentials in Aqueous Solutions; Marcel Dekker: New York, 1985. (21) Dhar, R. K.; Zheng, Y.; Rubenstone, J.; van Geen, A. A rapid colorimetric method of measuring arsenic concentrations in groundwater. Anal. Chim. Acta 2004, 526, 203−209. (22) Huang, K.-C.; Couttenye, R. A.; Hoag, G. E. Kinetics of heat assisted persulfate oxidation of methyl tert-butyl ether (MTBE). Chemosphere 2002, 49, 413−420. (23) Wang, Z.; Bush, R. T.; Liu, J. Arsenic(III) and Iron(II) cooxidation by oxygen and hydrogen peroxide: divergent reactions in the presence of organic ligands. Chemosphere 2013, 93, 1936−1941. (24) Sánchez-España, J.; López-Pamo, E.; Santofimia, E.; Aduvire, O.; Reyes, J.; Barettino, D. Acid mine drainage in the Iberian Pyrite Belt (Odiel river watershed, Huelva, Spain): Geochemistry, Mineralogy and Environmental implications. Appl. Geochem. 2005, 20, 1320−1356. (25) Rastogi, A.; Al-Abed, S. R.; Dionysiou, D. D. Sulfate radicalbased ferrous−peroxymonosulfate oxidative system for PCBs degradation in aqueous and sediment systems. Appl. Catal., B 2009, 85, 171−179. (26) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. Critical review of rate constants for reactions of hydrated electrons, hydrogen atoms and hydroxyl radicals in aqueous solutions. J. Phys. Chem. Ref. Data 1988, 17, 513−886. (27) Selim Reza, A. H. M.; Jean, J.-S.; Lee, M.-K.; Liu, C.-C.; Bundschuh, J.; Yang, H.-J.; Lee, J.-F.; Lee, Y.-C. Implications of organic matter on arsenic mobilization into groundwater: Evidence from northwestern (Chapai-Nawabganj), central (Manikganj) and southeastern (Chandpur) Bangladesh. Water Res. 2010, 44, 5556−5574. (28) Giasuddin, A. B. M.; Kanel, S. R.; Choi, H. Adsorption of humic acid onto nanoscale zerovalent iron and its effect on arsenic removal. Environ. Sci. Technol. 2007, 41, 2022−2027. (29) Lin, C.; Lin, K.-S. Photocatalytic oxidation of toxic organohalides with TiO2/UV: the effects of humic substances and organic mixtures. Chemosphere 2007, 66, 1872−1877.
ASSOCIATED CONTENT
* Supporting Information S
Experimental method details and additional experimental evidence. This material is available free of charge via the Internet at http://pubs.acs.org.
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AUTHOR INFORMATION
Corresponding Authors
*E-mail:
[email protected]. *E-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS This work was financially supported by the Australian Research Council Linkage Grant Scheme (Grant No. LP110100732 “Electron flow in iron hyper-enriched acidifying coastal environments”), National Natural Science Foundation of China (NSFC; Grant Nos. 21377023, 41273108), CRC CARE Project 6.6.01.06/07, and “Chen Guang” project (10CG34).
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REFERENCES
(1) Khan, A. W.; Ahmad, S. A.; Sayed, M. H. S. U. Arsenic contamination in groundwater and its effect on human health with particular reference to Bangladesh. J. Prev. Soc. Med. 1997, 16, 65−73. (2) Cullen, W. R.; Reimer, K. J. Arsenic speciation in the environment. Chem. Rev. 1989, 89, 713−764. (3) Borho, M.; Wilderer, P. Optimized removal of arsenate(III) by adaptation of oxidation and precipitation processes to the filtration step. Water Sci. Technol. 1996, 34, 25−31. (4) Pettine, M.; Campanella, L.; Millero, F. J. Arsenite oxidation by H2O2 in aqueous solutions. Geochim. Cosmochim. Acta 1999, 63, 2727−2735. (5) Hug, S. J.; Leupin, O. Iron-catalyzed oxidation of arsenic(III) by oxygen and by hydrogen peroxide: pH-dependent formation of oxidants in the Fenton reaction. Environ. Sci. Technol. 2003, 37, 2734− 2742. (6) Katsoyiannis, I. A.; Ruettimann, T.; Hug, S. J. pH dependence of Fenton reagent generation and As(III) oxidation and removal by corrosion of zero valent iron in aerated water. Environ. Sci. Technol. 2008, 42, 7424−7430. (7) Pang, S.-Y.; Jiang, J.; Ma, J. Oxidation of sulfoxides and arsenic(III) in corrosion of nanoscale zero valent iron by oxygen: Evidence against ferryl ions (Fe(IV)) as active intermediates in Fenton reaction. Environ. Sci. Technol. 2011, 45, 307−312. (8) Lee, H.; Choi, W. Photocatalytic oxidation of arsenite in TiO2 suspension: Kinetics and mechanisms. Environ. Sci. Technol. 2002, 36, 3872−3878. (9) Pignatello, J. J.; Oliveros, E.; MacKay, A. Advanced oxidation processes for organic contaminant destruction based on the Fenton reaction and related chemistry. Crit. Rev. Environ. Sci. Technol. 2006, 36, 1−84. 3984
dx.doi.org/10.1021/es405143u | Environ. Sci. Technol. 2014, 48, 3978−3985
Environmental Science & Technology
Article
(30) Ingham, M.; McConchie, J. A.; Wilson, S. R.; Cozens, N. Measuring and monitoring saltwater intrusion in shallow unconfined coastal aquifers using direct current resistivity traverses. J. Hydrol. 2006, 45, 69−82. (31) Lou, X.; Guo, Y.; Xiao, D.; Wang, Z.; Liu, J. Rapid dye degradation with reactive oxidants generated by chloride-induced peroxymonosulfate activation. Environ. Sci. Pollut. Res. 2013, 20, 6317− 6323. (32) Dodd, M. C.; Vu, N. D.; Ammann, A.; Le, V. C.; Kissner, R.; Pham, H. V.; Cao, T. H.; Berg, M.; von Gunten, U. Kinetics and mechanistic aspects of As(III) oxidation by aqueous chlorine, chloramines, and ozone: relevance to drinking water treatment. Environ. Sci. Technol. 2006, 40, 3285−3292. (33) Lente, G.; Kalmár, J.; Baranyai, Z.; Kun, A.; Kék, I.; Bajusz, D.; Takács, M.; Veres, L.; Fábián, I. One- versus two-electron oxidation with peroxomonosulfate ion: reactions with iron(II), vanadium(IV), halide ions, and photoreaction with cerium(III). Inorg. Chem. 2009, 48, 1763−1773. (34) Ball, D. L.; Edwards, J. O. The Kinetics and mechanism of the decomposition of caro’s acid. I. J. Am. Chem. Soc. 1956, 78, 1125− 1129. (35) Neppolian, B.; Doronila, A.; Grieser, F.; Ashokkumar, M. Simple and efficient sonochemical method for the oxidation of arsenic(III) to arsenic(V). Environ. Sci. Technol. 2009, 43, 6793−6798.
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