Activation of Peroxymonosulfate by Surface-Loaded Noble Metal

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Activation of Peroxymonosulfate by Surface-Loaded Noble Metal Nanoparticles for Oxidative Degradation of Organic Compounds Yong-Yoon Ahn, Eun-Tae Yun, Jiwon Seo, Changha Lee, Sang Hoon Kim, Jae-Hong Kim, and Jaesang Lee Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.6b02841 • Publication Date (Web): 26 Aug 2016 Downloaded from http://pubs.acs.org on September 8, 2016

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Environmental Science & Technology

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Activation of peroxymonosulfate by surface-loaded noble metal

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nanoparticles for oxidative degradation of organic compounds

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Yong-Yoon Ahn1, Eun-Tae Yun1, Ji-Won Seo2, Changha Lee2, Sang Hoon Kim3, Jae-Hong Kim4,

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and Jaesang Lee1,5*

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1

Civil, Environmental, and Architectural Engineering, Korea University, Seoul 136-701, Korea

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Urban and Environmental Engineering, KIST-UNIST-Ulsan Center for Convergent Materials (KUUC), Ulsan National Institute of Science and Technology, Ulsan 698-805, Korea 3

Center for Materials Architecturing, Korea Institute of Science and Technology (KIST), Seoul 136-701, Korea

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Chemical and Environmental Engineering, Yale University, New Haven, Connecticut 06511, United States

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Energy Environmental Policy and Technology, Green School, Korea University-KIST, Seoul 136-701, Korea

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*Corresponding author: E-mail: [email protected]; phone: +82-2-3290-4864; fax: +82-2-928-7656

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Abstract. This study demonstrates the capability of noble metal nanoparticles immobilized on

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Al2O3 or TiO2 support to effectively activate peroxymonosulfate (PMS) and degrade select

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organic compounds in water. The noble metals outperformed a benchmark PMS activator such as

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Co2+ (water-soluble) for PMS activation and organic compound degradation at acidic pH and

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showed the comparable activation capacity at neutral pH. The efficiency was found to depend on

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the type of noble metal (following the order of Pd > Pt ≈ Au >> Ag), the amount of noble metal

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deposited onto the support, solution pH, and the type of target organic substrate. In contrast to

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common PMS-activated oxidation processes that involve sulfate radical as a main oxidant, the

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organic compound degradation kinetics were not affected by sulfate radical scavengers and

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exhibited substrate dependency that resembled the PMS activated by carbon nanotubes. The

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results presented herein suggest that noble metals can mediate electron transfer from organic

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compounds to PMS to achieve persulfate-driven oxidation, rather than through reductive

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conversion of PMS to reactive sulfate radical.

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Keywords: noble metal, peroxymonosulfate activation, oxidative degradation, sulfate radical,

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electron transfer mediator

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INTRODUCTION

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The persulfate-based oxidation (persulfate herein represents ions or compounds

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containing SO52− or S2O82−.1) has been increasingly recognized as a viable, alternative oxidation

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process for water treatment and soil remediation.2-4 The oxidation is carried out as relatively

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stable persulfates such as peroxymonosulfate (HSO5−, PMS) and peroxydisulfate (S2O82−, PDS)

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are activated on site to generate reactive species such as a sulfate radical (SO4•−). Various

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strategies that can be readily put into practice have been developed to effectively activate

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persulfate through energy or electron transfer reactions. For instance, UV photolysis2, 5, 6 and

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thermolysis7 activate PMS and PDS by initiating intramolecular electron transfer and the

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associated homolytic cleavage of peroxide bonds within the persulfate molecule. Alternatively,

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persulfates can be electrochemically activated, for example, through cathodic reduction in an

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electrolytic cell.8 The activation through the similar reductive pathway can be also achieved

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using transition metals (e.g., Fe2+ and Co2+) and their elemental or oxide counterparts (e.g., Fe0,

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Fe(OH)3, β-MnO2, Mn2O3, and Co3O4) that reduce persulfates concomitantly with their facile

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oxidation.9-13 Conduction band electrons in semiconductors generated via photocatalysis have

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been also exploited to reduce persulfate.14

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A typical persulfate-activated oxidation process relies on sulfate radical’s strong

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oxidizing power (E0(SO4•−/SO42−) = 2.43 VNHE15) toward a broad spectrum of recalcitrant

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pollutants.2-4, 8, 16 While sulfate radical is less reactive than hydroxyl radical (•OH),17, 18 previous

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studies suggested a superior performance compared to well-established H2O2-based advanced

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oxidation for the degradation of select organic compounds.7, 19, 20 Unlike H2O2 activation that can


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be achieved with only a few transition metal ions (i.e., Fenton and Fenton-like reactions) in

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acidic condition or at circumneutral pH only when metal ions are coordinated to organic

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ligands21 or metal oxide surfaces,22 various activation strategies that are effective over a wider

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range of pH have been reported for the persulfate system. The oxidation by SO4•− favors direct

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electron abstraction and can even transform some anions (e.g., Cl− and OH−) into the

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corresponding radicals, which is not significant with •OH-induced oxidation.5 In particular,

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chlorine radical (Cl•) and active chlorine species (i.e., Cl2, HOCl) formed via one-electron

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oxidation of Cl− kinetically enhanced treatment of organic compounds and caused the occurrence

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of chlorinated intermediates.23, 24

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Recent studies16,

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have proposed that some persulfate activation schemes do not

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involve radical formation, contrasting the persulfate-based oxidation from classical advanced

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oxidation processes (AOPs). In a hypothesized non-radical mechanism postulated in these

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studies, persulfate is thought to function as an oxidant that directly accept electrons from an

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organic substrate, i.e., electron donor, with the help of an electron transfer mediator. Hence, the

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electron transfer mediator activates the organic compound degradation just like abovementioned

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persulfate activators (e.g., Co2+), but through a different mechanism. While the direct evidence

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for electron transfer through the activator is currently lacking in literature, many experimental

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observations indirectly suggest a high likelihood of this pathway occurring in a few processes.

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For example, the kinetics of 2,4-dichlorophenol oxidation by CuO/PDS (i.e., CuO as an

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activator) was found to be unaffected even when excess amounts of ethanol and chloride ions as

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a SO4•− scavenger were added.26 When carbon nanotubes (CNTs) were used to activate

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persulfate and degrade organic compounds, no radicals were detected (e.g., through electron spin



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resonance) and the degradation kinetics was also not retarded by excess radical scavenger.16 The

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mechanism involving electron transfer mediation by CNTs is also consistent with the observation

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that PMS activation was enhanced by nitrogen doping in CNT, the common approach to enhance

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conductivity of CNT.25

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We herein explore the potential use of noble metal nanoparticles as a PMS activator for

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the first time in literature, motivated by the recent emergence of persulfate-based oxidation

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schemes that involve electron transfer mediation mechanism. Noble metals are of our interest

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since they exhibit not only high electrical conductivity27 but also excellent capability to induce

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catalytic reduction28,

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mechanisms, persulfate reduction to SO4•− versus electron transfer from organics to persulfate).

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We examined whether most common noble metal catalysts such as platinum (Pt), palladium

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(Pd), gold (Au), and silver (Ag) supported on metal oxides (referred to in this study as NM-

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Al2O3 and NM-TiO2) can activate PMS and consequently degrade select organic pollutants. The

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effects of reaction parameters such as the catalyst loading and initial pH were investigated, and

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the dependence of the PMS activation efficiency on the substrate type was evaluated. In an effort

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to address the critical knowledge gap regarding the primary activation mechanism, we compared

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three PMS activating systems involving Pd-Al2O3, Co2+, and MWCNTs in terms of 1) the

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substrate specific nature of the oxidizing capacity and 2) the extent of reduction in performance

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in the presence of methanol, a radical scavenger. Finally, we assessed how NM-Al2O3 and NM-

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TiO2 would perform when used repeatedly.

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(as aforementioned, persulfate activation is initiated via two distinct

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MATERIALS AND METHODS


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Chemicals and Materials. The chemicals that were used as-received in this study include:

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aluminum oxide (Al2O3, PURALOX TH 100/150, Sasol; BET surface area = 150 m2/g; average

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particle size = 35 µm), titanium dioxide (TiO2, DT-51, Cristal Global; BET surface area = 90

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m2/g; average particle size = 1.7 µm), cobalt nitrate hexahydrate (Sigma-Aldrich), multiwall

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carbon nanotubes (MWCNTs, Hanwha), potassium monopersulfate (OXONE®, Sigma-Aldrich),

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potassium peroxydisulfate (Sigma-Aldrich), benzoic acid (Sigma-Aldrich), bisphenol A

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(Aldrich), carbamazepine (Sigma-Aldrich), 4-chlorophenol (Aldrich), 4-nitrophenol (Aldrich),

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phenol (Sigma-Aldrich), 2,4,6-trichlorophenol (Aldrich), pentachlorophenol (Aldrich), methanol

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(J.T. Baker), potassium iodide (Sigma-Aldrich), perchloric acid (Sigma-Aldrich), sodium

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bicarbonate (Sigma-Aldrich), sodium hydroxide (Fluka), phosphoric acid (Aldrich), and

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acetonitrile (J.T. Baker). All chemicals used in this study were of the highest purity available,

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and used without further treatment. Ultrapure water (>18 MΩ•cm) produced by a Milli-Q Water

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Purification System (Millipore) was used to prepare all the solutions.

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Catalyst Preparation and Characterization. Noble metals were deposited onto the surface of

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Al2O3 or TiO2 support using a coaxial pulsed arc-plasma deposition (APD) system (ULVAC,

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ARL-300) equipped with a rod-shaped noble metal cathode, a trigger electrode at the center of a

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reaction chamber, and a cylindrical anode coaxially aligned to surround the cathode.30 Metal

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oxide powders (Al2O3 or TiO2) were placed at the bottom of the chamber, directed toward the

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plasma source, and constantly stirred for effective dispersion. An electric charge was

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accumulated in a discharge condenser (capacity = 1080 µF; connected to the cathode) at the

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discharge voltage of 200 V. An ionized metal plasma was then instantly produced at room



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temperature under vacuum (10-5 Torr) by a trigger pulse on the cathode which subsequently

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formed nanoparticles on the surface of Al2O3 or TiO2.

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Energy dispersive spectroscopy (EDS) analysis (Figures S1-4) confirmed successful

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surface loading of noble metals on Al2O3. The transmission electron microscopy (TEM;

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TECNAI G2 F30ST, JEOL Ltd.) images in Figure 1 (Al2O3) and Figure S5 (TiO2) show that

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noble metal nanoparticles with an average size of approximately 2–5 nm were uniformly

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deposited on the surface of metal oxide supports, though a minor fraction of Au and Ag particles

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were observed to undergo aggregation (Pd (on Al2O3) = 3.10 ± 0.965 nm; Pt (on Al2O3) = 3.02 ±

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0.735 nm; Au (on Al2O3) = 3.35 ± 0.933 nm; Ag (on Al2O3) = 2.09 ± 1.124 nm; Pd (on TiO2) =

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2.76 ± 0.576 nm; Pt (on TiO2) = 2.02 ± 0.480 nm). The loading of Pd, Pt, Au, and Ag on both

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Al2O3 and TiO2 surface was estimated to be approximately 1 wt% using inductively coupled

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plasma mass spectrometry (ICP-MS, ELAN DRC-II, Perkin Elmer). The oxidation states of the

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noble metals deposited on Al2O3 were analyzed by X-ray photoelectron spectroscopy (XPS, PHI

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X-tool, ULVAC-PHI, Inc.) using the Al Kα line (1486.7 eV) as an excitation source. The

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spectrum of freshly-prepared Pd-Al2O3, which was essentially the same when TiO2 was an

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alternative support, is characterized by a double peak centered at 335.1 eV for Pd 3d5/2 and

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340.35 eV for Pd 3d3/2, suggesting the presence of Pd in a metallic form (Figure S6a).31 The Pt 4f

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band in the XPS spectrum of Pt-Al2O3 deconvolutes into three peaks at 74.3, 72.5, and 71.2 eV

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that correspond to Pt(IV), Pt(II), and Pt(0), respectively, indicating the co-existence of metallic

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and oxidized Pt (Figure S6b).32 The XPS spectrum of Au deposited on Al2O3 (Figure S6c) shows

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the Au 4f peaks at a binding energy of 83.8 for Au 4f7/2 and 87.5 eV for Au 4f5/2; these peaks are

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distinctive for Au(0).33 The XPS signal of Ag consisted of a doublet with two peaks at 368 and


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374 eV, respectively assigned to Ag 3d5/2 and Ag 3d3/2 (Figure S6d); this signal is attributed to

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surface deposition of metallic Ag.34 The immobilized noble metal particles were stored in the

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sealed borosilicate glass vials in the dark. PMS activation capacity of NM-Al2O3 (or NM-TiO2)

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samples was confirmed to remain unchanged after 2 years of storage. The zeta potentials of the

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aqueous-suspended NM-Al2O3 were measured as a function of pH using an electrophoretic light

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scattering spectrophotometer (ELSZ-1000, Otsuka Electronics).

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To explore the possible leaching of noble metals from NM-Al2O3, we prepared aqueous

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suspensions containing 0.25 g/L NM-Al2O3, 0.25 mM PMS and 0.1 mM 4-CP (4-chlorophenol;

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initially adjusted to pH 3, pH 7, and pH 10), performed PMS activation, and quantified the

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dissolved metal ions in the filtrates using ICP-MS. Chemical stability of noble metal-based

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activators was also examined based on determination of the leached metal ions during multiple

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catalytic cycles.

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Experimental Procedure and Analytical Methods. Degradation of various target organic

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compounds (0.1 mM) was monitored in a magnetically-stirred 40-mL reactor containing air-

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equilibrated aqueous suspensions of 0.25 g/L NM-Al2O3 (or NM-TiO2) and 0.25 mM PMS. The

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reaction suspensions were typically buffered at ca. 7.0 using 1 mM bicarbonate buffer and found

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to change only marginally over the course of reaction. We confirmed that bicarbonate could

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neither kinetically affect PMS activation processes (Figure S7) nor activate PMS by itself

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(Figure S8) (note that some anions (e.g., HPO42−, Cl−, HCO3−) were demonstrated to active

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PMS7 and carbonate/bicarbonate can react with SO4•− to produce some radical species (e.g.,

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CO3•−).35 For select experiments performed to evaluate pH effect, the suspensions were



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unbuffered and the initial pH was adjusted using 0.1 M HClO4 or NaOH solution. Sample

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aliquots were periodically withdrawn from the reactor using a 1 mL syringe, filtered through a

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0.45 µm PTFE filter (Millipore), and injected into a 2 mL amber glass vial containing excess

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methanol (0.5 M) to quench any residual radicals. The concentration of organic compounds were

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determined using a HPLC (Agilent Infinity 1260) system equipped with a C-18 column

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(ZORBAX Eclipse XDB-C18) and a UV/Vis detector (G1314F 1260VWD). The typical eluent

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consisted of a binary mixture of 0.1% (v/v) aqueous phosphoric acid solution and acetonitrile

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(45/55 by volume). Intermediates formed during 4-CP oxidation by the activated PMS were

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qualitatively identified using the Rapid Separation Liquid Chromatography (RSLC)/orbitrap

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MS/MS system. The RSLC separation was carried out on an acclaimTM C18 column (150 mm ×

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2.1 mm, 2.2 µm; Thermo Fisher Scientific Inc.) with a mobile phase consisting of 0.1 % aqueous

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formic acid solution and acetonitrile. The mass analysis was performed in a negative electrospray

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ionization (ESI) mode. The relative abundances of oxidation intermediates were estimated based

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on their mass spectroscopy peak areas. Accurate mass measurements were guaranteed with the

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low ppm range (< 10 ppm of the theoretical mass). Evolution of chloride ions as a result of

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dechlorination of 4-chlorophenol was monitored using an ion chromatography (IC, Dionex DX-

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120) equipped with a Dionex IonPac AS-14 and a conductivity detector. PMS was

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spectrophotometrically quantified using the method proposed by Liang et al., based on an iodine

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(λmax = 352 nm) forming reaction between PMS and iodide.36 A calibration plot of absorbance at

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352 nm versus PMS concentration showed a linear relationship over the concentration range of

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30 µM to 500 µM, suggesting 30 µM as the detection limit for PMS. Formaldehyde that forms as

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a result of methanol oxidation was quantitatively measured using the HPLC after derivatization

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with 2,4,-dinitrophenylhydrazine (DNPH).37 Gaseous CO2 production as an evidence for


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mineralization of organic compounds was monitored using a gas chromatography (GC, Agilent

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HP6890A) that was equipped with a flame ionization detector (FID), a Carboxen 1010 PLOT

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capillary column, and a CO2 methanizer (Agilent G2474A). Dissolved organic carbon was

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measured using a total organic carbon analyzer (TOC-VCPH, Shimadzu). For electron

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paramagnetic resonance (EPR) analysis, 5,5-dimethyl-1-pyrroline-N-oxide (DMPO) was used as

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a spin-trapping agent for SO4•−. The EPR spectra were monitored in the Pd-Al2O3/PMS system

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using a JES-TE 300 spectrometer (JEOL, Japan) under the following conditions: microwave

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power = 1 mW, microwave frequency = 9.421 GHz, center field = 3375 G, modulation width =

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0.2 mT, and modulation frequency = 100 kHz.

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RESULTS AND DISCUSSION

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PMS Activation by Noble Metals. Results in Figure 2 demonstrate that PMS effectively

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degraded the model organic compound, 4-CP, but only when Pd-Al2O3 or Pd-TiO2 was present.

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Neither Pd-Al2O3, Pd-TiO2, nor PMS alone caused any decrease in 4-CP concentration (Figure

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S8). PMS also did not affect 4-CP concentration at all when bare Al2O3 or TiO2 without noble

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metals was used (Figure S8). The oxidation of 4-CP by the Pd-Al2O3 or Pd-TiO2 was

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accompanied by the loss of PMS that is related to the stoichiometric reduction into sulfate ions

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(SO42−) (Figure 2). Pd-Al2O3 (or Pd-TiO2) decomposed 4-CP in the presence of PMS at a rate

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comparable to Co2+, a benchmark PMS activator,3 while PMS consumption was 2.5 times faster

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with Co2+ than with Pd-Al2O3 (Figure S9). The rapid PMS depletion in the Co2+/PMS system is

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ascribed to oxidative conversion of PMS by Co3+ and concurrent regeneration of Co2+ (HSO5− +

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Co3+ → HSO5• + Co2+).38 Approximately 33.7 % of chlorine initially present in 4-CP was found



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to be released as chloride ion (Figure 2), which is consistent with 38.2 % dechlorination of 4-CP

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by Co2+/PMS system (Figure S9). The comparable 4-CP degradation between Pd and Co2+ is

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noteworthy, considering that Pd is in an immobilized, solid form in contrast to soluble Co2+. The

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gaseous CO2 produced after 8 h of 4-CP oxidation by either Pd-Al2O3/PMS or Co2+/PMS

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corresponded to only ca. 1.2-1.3 % mineralization of the initial 4-CP concentration (Figure S10).

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Other noble metals also induced PMS activation and 4-CP degradation when deposited

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onto Al2O3 (a more inert, non-reducible support used in this set of experiments to isolate the

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noble metal effects39), but to a widely varying degrees depending on both the type and amount of

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noble metal (Figure 3 and Figure S11). Over the range of metal loadings investigated, Pd was

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found most effective in 4-CP degradation. While a greater difference was observed at lower

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metal loadings, on average, the initial 4-CP degradation rate was 31% faster with Pd compared to

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the second most effective Au. No significant difference in the 4-CP degradation kinetics was

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found between Au and Pt. Unlike other metals, no reduction in the 4-CP concentration was

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observed when Ag was used, similar to the report that silver ion cannot activate PMS and

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therefore is ineffective in chlorophenol degradation.9

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The reason behind the effect of metal type is unknown, while the efficient PMS

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activation by Pd-, Au-, and Pt-Al2O3 might be related to the intrinsic surface catalytic property of

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these metals, for example, in reduction of nitrophenol29, 40 and nitrate.28 When Pd was used as an

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activator, the reaction rate reaches near plateau values at the 0.375 g/L, above which further

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loading induced only a marginal improvement in 4-CP removal kinetics. On the other hand, 4-CP

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degradation efficiency of Au- and Pt-Al2O3 continued to increase in proportion to the loading


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amount, becoming comparable to that of Pd-Al2O3 at 0.5 g/L loading. The PMS decomposition

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showed a similar trend as 4-CP degradation (Figure 3 and Figure S12), with the decay rate

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following the same order of Pd > Au ≈ Pt >> Ag and increasing with greater metal loading.

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These results collectively suggest that select noble metals, when deposited onto metal oxide

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surface, can activate PMS to result in PMS reduction and subsequent degradation of organic

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compound.

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The Effect of Initial pH on PMS Activation Efficiency. The kinetics of 4-CP degradation and

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PMS reduction were not significantly affected by the change in solution pH between 3 to 9

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(Figure 4 and Figure S13), when PMS was activated by Pd-Al2O3 or Pt-Al2O3. When the

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experimental suspensions were initially adjusted to an acidic or a basic pH and were buffered at

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pH 7 using 1 mM bicarbonate, the change in solution pH was marginal over the course of PMS

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activation. This relative pH independence, as confirmed in other persulfate-based oxidation

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processes, is one notable advantage compared to other peroxide-based oxidation processes such

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as H2O2/Fe2+ the efficiency of which is highly pH sensitive. When pH was raised to 11, however,

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4-CP degradation kinetics was significantly inhibited in both cases. These results are in contrast

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to the pH dependence of 4-CP oxidation with Au-Al2O3, in which alkaline and neutral pH

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favored but acidic pH retarded 4-CP degradation. Almost no 4-CP decomposition was observed

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with Ag-Al2O3, regardless of the initial pH. These 4-CP degradation kinetics correlated generally

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well with PMS decomposition rates for all the composite catalysts over all pH range investigated

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(Figure 4 and Figure S13). For example, when Pd-Al2O3 or Pt-Al2O3 was used as a PMS

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activator, the efficient 4-CP degradation was accompanied by the fast PMS decomposition

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between pH 3 and 9 and the slow 4-CP degradation by the slow PMS decomposition at pH 11.



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Likewise, an increase in pH accelerated both 4-CP degradation and PMS decomposition in Au-

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Al2O3 case. The inability of Ag-Al2O3 to degrade 4-CP in the presence of PMS accords with no

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detectable decomposition of PMS in the pH range investigated.

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The quantitative monitoring of noble metal leaching from NM-Al2O3 under varying pH

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conditions (Table S1) confirms negligible loss of noble metal mass (less than 0.03 % of the

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initial mass) during PMS activation, regardless of initial pH. This rules out the possibility that

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the pH-dependent PMS activation efficiency of NM-Al2O3 (Figure 4) is attributable to the

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chemical stability of activators that may be sensitive to pH conditions. The decreased efficiency

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of 4-CP degradation by Pd (or Pt)-Al2O3 at pH 11 could be partly related to the change of

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activator particle surface charge from positive to negative; the PZC (point of zero charge) of

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activator particle was found to be at around 10 (Figure S14), while PMS exists as anion over the

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entire pH range (pKa1 = 0.4 and pKa2 = 9.3).41, 42 We exclude the onset of electrostatic repulsion

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between negatively charged activator surface and deprotonated 4-CP as the major cause of

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decreased kinetics at pH 11. We found that when Pd-Al2O3 was employed as an activator, other

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chlorophenols, 2,4,6-trichlorophenol (2,4,6-TCP) and pentachlorophenol (PCP) also underwent

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rapid degradation in the pH range of 3 to 9 and showed drastically retarded decomposition

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kinetics only at pH 11, despite widely varying pKa values (i.e., pKa(4-CP) = 9.4143; pKa(2,4,6-

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TCP) = 6.2343; pKa(PCP) = 4.7044) (Figure S15). Alternatively, hydrolysis of PMS under

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alkaline conditions (SO52− + H2O → HO2− + SO42− + H+)45 might be responsible for futile loss of

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PMS. However, we found that PMS does not decay over the entire pH range in the absence of

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noble metal (but in the presence of bare Al2O3). When Pd-Al2O3 was added alternatively, PMS

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decayed over the entire pH range except only pH 11 (Figure S16), further excluding this


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hypothesis. These collectively suggest that the ineffective PMS activation by Pd (Pt)-Al2O3 is a

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major cause of the decelerated 4-CP degradation at high pH, possibly due to electrostatic

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repulsion between NM and PMS or through a mechanism currently unknown (e.g., the formation

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of poisoning oxide layer on Pd (Pt) surface at strongly alkaline pH). It is noteworthy that Au that

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underwent negligible reduction in PMS activation capacity at pH 11 (Figure 4) is much less

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prone to surface oxidation than Pd and Pt and is superior over Pt and Pd in (electro)catalytic O2

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reduction and alcohol oxidation in alkaline media.46, 47 Considering its relative irrelevance to

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realistic conditions, we did not further delve into the phenomenon occurring at such an extreme

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high pH condition.

288 289

The above pH dependence of 4-CP degradation kinetics presents a unique feature of PMS

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activation by NM-Al2O3 compared to benchmark Co2+ activator (Figure S15). The Co2+ showed

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the optimal 4-CP degradation in the near-neutral pH range, due to predominance of CoOH+

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species in the neutral pH region which is known to be an efficient PMS activator48 and

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precipitation of cobalt ions under even slightly alkaline conditions.49, 50 At low pH, the loss of

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oxidation capacity has been attributed to the decrease of oxidizing power of SO4•−,49 a main

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radical oxidant in typical persulfate-based oxidation processes such as the one employing Co2+ as

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an activator. It is therefore interesting to note that chlorophenols were effectively degraded when

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PMS was activated by Pd-Al2O3 and Pt-Al2O3 at pHs as low as 3.0 (Figure 4 and Figure S15).

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This not only highlights the advantage of using NM-Al2O3 as an activator for PMS but indicates

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likely involvement of a different persulfate reduction mechanism that does not result in SO4•−.

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Degradation of Various Organics. The efficacy of PMS activated by NMs, deposited onto

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either Al2O3 or TiO2, to degrade various organic compounds was further evaluated (Figure 5).

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The degradation rate was consistently in the order of Pd > Pt ≈ Au >> Ag, regardless of the type

304

of target organic compounds and metal oxide support, while some variations were observed. In

305

general, TiO2 support caused faster organic compound degradation for the same NM loading

306

under acidic and neutral conditions (Figure S17 and Figure 5). This effect was more pronounced

307

for phenolic compounds such as phenol (PH) and bisphenol A (BPA), presumably due to binding

308

of phenolic hydroxyl group with surface titanol moiety (>TiOH).51 Most notable in Figure 5 is

309

the dependence of degradation rate on the type of organic compounds. PH and chlorophenols (4-

310

CP and 2,4,6-TCP) were readily oxidized, while other compounds were degraded only

311

moderately (BPA and carbamazepine (CBZ)) or almost negligibly (benzoic acid (BZA) and 4-

312

nitrophenol (4-NP)). This feature once again appears unique to NM-Al2O3 (or NM-TiO2)/PMS

313

system, since PMS activated through other routes has been claimed to be effective for a wide

314

range of organic compound degradation.3, 8, 9, 52 For example, our previous study also suggested

315

that, through activation with zero-valent iron, PMS forms SO4•− that effectively oxidized various

316

organic compounds following rather constant kinetics rates.52 Selective nature toward organic

317

degradation observed in Figure 5 once again suggests that activation of PMS by NM-Al2O3 (or

318

NM-TiO2)/PMS and resulting organic compound oxidation might not involve SO4•− as a primary

319

oxidant.

320 321

PMS Activation Mechanism. To confirm the above speculation, excess methanol was added as

322

a SO4•− quencher and the kinetics of 4-CP oxidation was evaluated (Figure 6a). If SO4•− is a

323

major oxidant that results from PMS activation, the addition of methanol would inhibit the


Page 15 of 34


324

organic compound degradation.52 Consistently, we observed that excess methanol completely

325

quenched 4-CP oxidation in the Co2+/PMS system, where PMS is known to be reductively

326

converted into SO4•−.3, 9 In contrast, the same concentration (i.e., 0.25 M) of methanol as used in

327

Co2+/PMS control experiment did not cause any noticeable retardation in the rate of 4-CP

328

degradation when Pd-Al2O3 and Au-Al2O3 were used for PMS activation. The experiments to

329

evaluate the conversion of methanol to formaldehyde (inset of Figure 6a) demonstrate that

330

Co2+/PMS initially caused a steep increase in formaldehyde concentration whereas methanol

331

oxidation very slowly proceeded or was absent within 15 min in the NM-Al2O3/PMS systems.

332

These results suggest that the role of SO4•− in the oxidation of organics by NM-Al2O3/PMS is

333

most likely negligible.

334 335

We further compared the oxidizing capacity of Pd-Al2O3/PMS to that of Co2+/PMS and

336

MWCNT/PMS in terms of substrate specificity (Figure 6b). Two activating systems, Pd-

337

Al2O3/PMS and MWCNT/PMS, showed the similar substrate specificity; i.e., organics that were

338

readily oxidized in Pd-Al2O3/PMS was also in general well oxidized in MWCNT/PMS. In

339

contrast, when Co2+ was employed to activate PMS, a significant difference in the substrate

340

specificity was observed. For example, CBZ that was recalcitrant to the degradation by both Pd-

341

Al2O3/PMS and MWCNT/PMS underwent the fastest decay with Co2+. The higher chlorinated

342

phenol underwent more rapid decomposition in the systems of Pd-Al2O3/PMS and

343

MWCNT/PMS, which is in contrast to the substrate-specificity of Co2+/PMS. It has been

344

reported that MWCNT activates PMS through a different mechanism in which electron is

345

transferred from the organics to PMS through CNTs that functions as an electron shuttle.16, 25 As

346

a result, organics are oxidized and PMS are reduced to sulfate. We believe that similar non-



Page 16 of 34

347

radical pathway is highly plausible for noble metals considering the materials’ superior electrical

348

conductivity.26 Accordingly, we did not observe any DMPO-SO4•− adduct in the EPR spectrum

349

of Pd-Al2O3/PMS (Figure S18). Instead, we observed the occurrence of peaks assigned to 5,5-

350

dimethylpyrrolidone-2-(oxy)-(1) (DMPOX) as a product of direct DMPO oxidation,53 which was

351

also observed when CNTs were alternative used for PMS activation.16 In contrast, the spectral

352

characteristics of DMPO-SO4•− adduct appeared in the Co2+/PMS system (Figure S18). There is

353

a likelihood that quinone intermediates (likely produced from 4-CP oxidation) may reductively

354

convert PMS to SO4•− or may mediate the transfer of electrons from 4-CP to PMS.54 However,

355

Figure S19 demonstrates that 4-CP degradation was not achieved when benzoquinone or

356

catechol was used as a PMS activator, which rules out a role of quinone intermediates in the

357

PMS activation mechanism.

358 359

The comparative analysis of reaction intermediates also confirmed the difference in PMS

360

activation mechanism between noble metals versus Co2+. For instance, hydroxybenzoquinone,

361

(2E,4Z)-3-hydroxyhexa-2,4-dienedioic acid, and 4-hydroxyphenylbenzoquinone appeared only

362

during oxidative degradation of 4-CP by Co2+/PMS (Table S3). In contrast, PMS activation by

363

Pd-Al2O3 led to formation of (Z)-4,5-dioxopent-2-enoic acid and some unidentified intermediates

364

that were not observed when Co2+/PMS decomposed 4-CP (Table 3). In particular, the time-

365

dependent profiles of the relative abundances of all detected compounds (Figure S20) showed

366

that the distribution of major intermediates varied depending on the type of activator: (Z)-4,5-

367

dioxopent-2-enoic acid and an unknown compound with m/z 186.11 formed as primary products

368

of 4-CP oxidation by Pd-Al2O3/PMS whereas they were not detectable in the Co2+/PMS system.

369

The quantification of quinone intermediates using HPLC (Figure S21) suggested that formation


Page 17 of 34


370

of hydroquinone (HQ) seemed more pronounced when Co2+ was used as an activator; almost

371

threefold higher concentration of HQ was detected in the Co2+/PMS system after 30 min of 4-CP

372

oxidation than in the Pd-Al2O3/PMS system.

373 374

If a role as a mediator in transferring electron from organics to PMS were critical to the

375

NM-induced PMS activation, PMS reduction would not occur without organic substances. Note

376

that PMS was reduced in the aqueous suspensions of CNTs only when organics were added.16

377

Figure S22 showed that PMS was not decomposed at all by Au-Al2O3, Pt-Al2O3, or CNTs in the

378

absence of 4-CP whereas 4-CP addition drastically decreased PMS concentration. This contrasts

379

with PMS degradation by Co2+ that was kinetically retarded but still significant in the absence of

380

4-CP (Figure S22). The acceleration of PMS reduction in the ternary mixture, Co2+/PMS/4-CP, is

381

likely attributable to PMS activation or Co2+ regeneration from Co3+ by quinone intermediates.55

382

Accordingly, the results confirm that NM-Al2O3 effectively facilitate electron transfer between

383

pollutants (electron donor) and PMS (electron acceptor), contributing to oxidative degradation

384

via non-radical mechanism. Unlike other tested activators, Pd-Al2O3 exhibited almost the same

385

but significant PMS degradation efficiency regardless of the presence of 4-CP. This may result

386

from the catalytic activity of Pd in chemical transformation of peroxide compounds.56 It is

387

noteworthy that Pd-Al2O3 repeatedly achieved PMS decomposition without significant loss of

388

catalytic activity when 4-CP was absent (Figure S23).

389 390

Repetitive Use. Figure 7 explores the possibility of using immobilized Pd and Pt more than one

391

time for PMS activation. Repetitive uses caused a gradual decline in the capability of the noble

392

metals to decompose 4-CP in the presence of PMS, but the extent of reduction of the 4-CP



Page 18 of 34

393

degradation rate (Figure 7) or the PMS reduction rate (Figure S24) considerably varied

394

depending on the noble metal and support material employed. Drastic deactivation of Pd-Al2O3

395

for oxidative degradation of 4-CP was observed after single use, with the removal rate being

396

decreased by 68 % only after the first cycle. Similarly, when Pt was used as the activator, the 4-

397

CP degradation efficiency was decreased by up to 80 % in the second cycle. When TiO2 was

398

used as an alternate host for the noble metals, the reusability of Pd and Pt in PMS activation was

399

significantly improved. For instance, the Pt-TiO2 samples recovered after the first and second

400

cycles were respectively still capable of oxidizing 47 % and 25 % of 4-CP within 60 min. But

401

Pd-TiO2/PMS was able to maintain most of its catalytic property for the first three cycles, while

402

subsequent reuse decreased the efficiency down to 75 % of its original oxidation capacity.

403 404

In the recycling experiments, each cycle was carried out in different batches using NM-

405

Al2O3 (or NM-TiO2) that was recovered after previous cycle by using a 0.45 µm PTFE

406

membrane filter, washed with 2 L distilled water, and resuspended in a freshly-prepared aqueous

407

solution of PMS. Therefore, it is not likely that the steady reduction in the 4-CP degradation

408

efficiency is caused by accumulation of oxidation products in the experimental suspensions. We

409

found that the morphological change of noble metals did not occur before and after PMS

410

activation (Figures S25, S26, and S27) and therefore would not be responsible for the gradual

411

reduction in PMS activation performance. We also did not detect any meaningful accumulation

412

of noble metals in the filtrates in trace amounts (less than 0.012 % of the initial mass) (Table S2).

413

The EDS analysis of five randomly-selected regions from fresh Pd-Al2O3 versus Pd-Al2O3

414

reused for five cycles showed marginal loss of Pd during the repeated uses in PMS activation

415

(i.e., 2.40 ± 1.76 atomic % for fresh Pd-Al2O3; 2.08 ± 1.03 atomic % for reused Pd-Al2O3). The


Page 19 of 34


416

comparison of their TEM images (Figure S28) further confirmed that Pd particles still remained

417

uniformly distributed on Al2O3 after five cycles. The results collectively exclude metal

418

dissolution as a primary cause. We suspect that reduced efficiency with repeated applications

419

may involve a change in the oxidation state of the noble metal, for example, through non-

420

conducting metal oxide formation. Note that PMS was reduced by Pd-Al2O3 (or Pd-TiO2) even

421

when 4-CP, an electron donor, was absent (Figures S8, S16, and S22), indicating potential

422

oxidation of Pd due to PMS. Also surface deposition of organic intermediates may hinder noble

423

metals from being available for PMS activation, resulting in a gradual performance reduction

424

during the catalytic use of NM-Al2O3 (or NM-TiO2). As opposed to steady retardation in PMS

425

degradation in the presence of 4-CP (Figure S24), the efficiency of Pd-Al2O3 for PMS

426

decomposition was constantly maintained even after the fifth cycle when 4-CP was absent

427

(Figure S23). Furthermore, significant TOC reduction (ca. 60 % mineralization) after 4-CP

428

oxidation by Pd-Al2O3/PMS (Figure S29) is contradictory to the negligible conversion of 4-CP to

429

CO2 (ca. 1.2 % mineralization) (Figure S10). Note that neither TOC reduction (Figure S29) nor

430

CO2 generation (Figure S10) could be achieved with the Co2+/PMS system. The results

431

collectively corroborate the accumulation of oxidation intermediates on the surface of noble

432

metal-based activators.

433 434

Environmental Applications. This study presents the first instance of successfully employing

435

noble metals for PMS activation. It is noteworthy that noble metal ions such as Pd2+, Pt4+, and

436

Ag+ were found ineffective for PMS decomposition (not PMS activation).57 Comparative

437

analysis of the activation process using various noble metal nanoparticles immobilized on Al2O3

438

and TiO2 suggested a few interesting features that contrast this strategy from existing ones. First,



Page 20 of 34

439

the Pd-Al2O3 exhibited comparable or even better efficiency compared to the conventional

440

activation system involving Co2+ in degrading phenolic compounds when the same amount of

441

activating reagent (i.e., 25 µM) was employed (Figure S30), while consuming much less PMS

442

(Figure S9). Second, Pd- and Pt-Al2O3 exhibited significant PMS activating capacity over a wide

443

pH range (pH 3 to 9), especially at low pH in which SO4•−-mediated oxidation does not

444

efficiently function.49 Third, contrary to the previously observed non-selective reactivity of

445

SO4•− toward phenolic compounds,52 the rate of oxidative decomposition by PMS activated with

446

NM-Al2O3 was found to be highly dependent on the target substrate. Particularly, the substrate-

447

specific reactivity likely allows the NM-Al2O3/PMS systems to maintain significant treatment

448

efficiency in the presence of background organic matter; we found that Pd-Al2O3/PMS was still

449

capable of effective 4-CP degradation in the presence of excess methanol or humic acid whereas

450

Co2+/PMS caused no 4-CP decomposition (Figures 6a and S31). Finally, the mechanism appears

451

to involve facile transfer of electrons from organics (electron donor) to PMS (electron acceptor)

452

through electron transfer mediation by NM. We expect that the same mechanism will function

453

when different persulfate species are used; for example, we found that significant degradation of

454

4-CP was also observed with Pd-Al2O3 (Figure S32) when PDS, instead of PMS, was used.

455

While this study introduces a new materials approach for persulfate-based oxidation process, it

456

also raises many questions that have not been addressed in literature and thus require further

457

studies, particularly on ways to effectively regenerate spent noble metal activators (given that

458

reduction in activation performance is attributed to the changes in their oxidation states). These

459

might include: 1) alternative use of other metal oxide supports (e.g., SiO2, CeO2, and Nb2O5) 2)

460

reduction using chemical reagents (e.g., NaBH4, H2 gas), and 3) photocatalytic treatment.58

461

Acknowledgements


Page 21 of 34


462

This research was supported by the Basic Science Research Program through the National

463

Research

464

(2014R1A1A2056935). This work was also partially supported by a grant from the future R&D

465

Program (2E26120) funded by Korea Institute of Science and Technology, by Korea Ministry of

466

Environment as “The GAIA Project” (2016000550007), and by the Ministry of Trade, Industry,

467

and Energy (MOTIE), Korea, as "Encouragement Program for the Industries of Economic

468

Cooperation Region (R0004881)".

Foundation

of

Korea

(NRF)

funded

by

the

Ministry

of

Education

469 470

Supporting Information Available.

471

This information is available free of charge via the Internet at http://pubs.acs.org/.

472 473

Literature Cited

474

1. Harald Jakob; Stefan Leininger; Thomas Lehmann; Sylvia Jacobi; Gutewort, S., Peroxo

475

Compounds, Inorganic. In Ullmann's Encyclopedia of Industrial Chemistry, John Wiley and

476

Sons: 2012; pp 293-324.

477

2. Liu, X. W.; Zhang, T. Q.; Zhou, Y. C.; Fang, L.; Shao, Y., Degradation of atenolol by

478

UV/peroxymonosulfate: Kinetics, effect of operational, parameters and mechanism.

479

Chemosphere 2013, 93, (11), 2717-2724.

480

3. Anipsitakis, G. P.; Dionysiou, D. D., Degradation of organic contaminants in water with

481

sulfate radicals generated by the conjunction of peroxymonosulfate with cobalt. Environ. Sci.

482

Technol. 2003, 37, (20), 4790-4797.

483

4. Rastogi, A.; Ai-Abed, S. R.; Dionysiou, D. D., Sulfate radical-based ferrous-

484

peroxymonosulfate oxidative system for PCBs degradation in aqueous and sediment systems.

485

Appl. Catal. B. Environ. 2009, 85, (3-4), 171-179.



486

5. Yang, Y.; Pignatello, J. J.; Ma, J.; Mitch, W. A., Comparison of halide impacts on the

487

efficiency of contaminant degradation by sulfate and hydroxyl radical-based advanced oxidation

488

processes (AOPs). Environ. Sci. Technol. 2014, 48, (4), 2344-2351.

489

6. Liu, X. W.; Fang, L.; Zhou, Y. C.; Zhang, T. Q.; Shao, Y., Comparison of UV/PDS and

490

UV/H2O2 processes for the degradation of atenolol in water. J. Environ. Sci. 2013, 25, (8), 1519-

491

1528.

492

7. Yang, S. Y.; Wang, P.; Yang, X.; Shan, L.; Zhang, W. Y.; Shao, X. T.; Niu, R., Degradation

493

efficiencies of azo dye Acid Orange 7 by the interaction of heat, UV and anions with common

494

oxidants: Persulfate, peroxymonosulfate and hydrogen peroxide. J. Hazard. Mater. 2010, 179,

495

(1-3), 552-558.

496

8. Chen, W. S.; Huang, C. P., Mineralization of aniline in aqueous solution by electrochemical

497

activation of persulfate. Chemosphere 2015, 125, 175-181.

498

9. Anipsitakis, G. P.; Dionysiou, D. D., Radical generation by the interaction of transition metals

499

with common oxidants. Environ. Sci. Technol. 2004, 38, (13), 3705-3712.

500

10. Liu, H. Z.; Bruton, T. A.; Doyle, F. M.; Sedlak, D. L., In situ chemical oxidation of

501

contaminated groundwater by persulfate: Decomposition by Fe(III)- and Mn(IV)-containing

502

oxides and aquifer materials. Environ. Sci. Technol. 2014, 48, (17), 10330-10336.

503

11. Oh, S. Y.; Kim, H. W.; Park, J. M.; Park, H. S.; Yoon, C., Oxidation of polyvinyl alcohol by

504

persulfate activated with heat, Fe2+, and zero-valent iron. J. Hazard. Mater. 2009, 168, (1), 346-

505

351.

506

12. Saputra, E.; Muhammad, S.; Sun, H. Q.; Ang, H. M.; Tade, M. O.; Wang, S. B., Manganese

507

oxides at different oxidation states for heterogeneous activation of peroxymonosulfate for phenol

508

degradation in aqueous solutions. Appl. Catal. B. Environ. 2013, 142, 729-735.

509

13. Yang, Q. J.; Choi, H.; Dionysiou, D. D., Nanocrystalline cobalt oxide immobilized on

510

titanium dioxide nanoparticles for the heterogeneous activation of peroxymonosulfate. Appl.

511

Catal. B. Environ. 2007, 74, (1-2), 170-178.

512

14. Kim, H.; Yoo, H. Y.; Hong, S.; Lee, S.; Lee, S.; Park, B. S.; Park, H.; Lee, C.; Lee, J.,

513

Effects of inorganic oxidants on kinetics and mechanisms of WO3 mediated photocatalytic

514

degradation. Appl. Catal. B. Environ. 2015, 162, 515-523.

515

15. Huie, R. E.; Clifton, C. L.; Neta, P., Electron transfer reaction rates and equilibria of the

516

carbonate and sulfate radical anions. Radiat. Phys. Chem. 1991, 38, (5), 477-481.


Page 22 of 34

Page 23 of 34


517

16. Lee, H.; Lee, H. J.; Jeong, J.; Lee, J.; Park, N. B.; Lee, C., Activation of persulfates by

518

carbon nanotubes: Oxidation of organic compounds by nonradical mechanism. Chem. Eng. J.

519

2015, 266, 28-33.

520

17. Zhou, D. N.; Zhang, H.; Chen, L., Sulfur-replaced Fenton systems: can sulfate radical

521

substitute hydroxyl radical for advanced oxidation technologies? J. Chem. Technol. Biotechnol.

522

2015, 90, (5), 775-779.

523

18. Toth, J. E.; Rickman, K. A.; Venter, A. R.; Kiddle, J. J.; Mezyk, S. P., Reaction kinetics and

524

efficiencies for the hydroxyl and sulfate radical based oxidation of artificial sweeteners in water.

525

J. Phys. Chem. A 2012, 116, (40), 9819-9824.

526

19. Yoon, S. H.; Jeong, S.; Lee, S., Oxidation of bisphenol A by UV/S2O82-: Comparison with

527

UV/H2O2. Environ. Technol. 2012, 33, (1), 123-128.

528

20. Shah, N. S.; He, X. X.; Khan, H. M.; Khan, J. A.; O'Shea, K. E.; Boccelli, D. L.; Dionysiou,

529

D. D., Efficient removal of endosulfan from aqueous solution by UV-C/peroxides: A

530

comparative study. J. Hazard. Mater. 2013, 263, 584-592.

531

21. Keenan, C. R.; Sedlak, D. L., Ligand-enhanced reactive oxidant generation by

532

nanoparticulate zero-valent iron and oxygen. Environ. Sci. Technol. 2008, 42, (18), 6936-6941.

533

22. Pham, A. L. T.; Doyle, F. M.; Sedlak, D. L., Kinetics and efficiency of H2O2 activation by

534

iron-containing minerals and aquifer materials. Wat. Res. 2012, 46, (19), 6454-6462.

535

23. Anipsitakis, G. P.; Dionysiou, D. D.; Gonzalez, M. A., Cobalt-mediated activation of

536

peroxymonosulfate and sulfate radical attack on phenolic compounds. Implications of chloride

537

ions. Environ. Sci. Technol. 2006, 40, (3), 1000-1007.

538

24. Anipsitakis, G. P.; Tufano, T. P.; Dionysiou, D. D., Chemical and microbial

539

decontamination of pool water using activated potassium peroxymonosulfate. Wat. Res. 2008,

540

42, (12), 2899-2910.

541

25. Duan, X. G.; Sun, H. Q.; Wang, Y. X.; Kang, J.; Wang, S. B., N-doping-induced nonradical

542

reaction on single-walled carbon nanotubes for catalytic phenol oxidation. ACS Catal. 2015, 5,

543

(2), 553-559.

544

26. Zhang, T.; Chen, Y.; Wang, Y. R.; Le Roux, J.; Yang, Y.; Croue, J. P., Efficient

545

peroxydisulfate activation process not relying on sulfate radical generation for water pollutant

546

degradation. Environ. Sci. Technol. 2014, 48, (10), 5868-5875.



547

27. Lee, J.; Kim, J.; Choi, W., TiO2 photocatalysis for the redox conversion of aquatic

548

pollutants. In Aquatic Redox Chemistry (ACS Symposium Series), Tratnyek, P. G.; Grundl, T. J.;

549

Haderlein, S. B., Eds. 2011; Vol. 1071, pp 199-222.

550

28. Gauthard, F.; Epron, F.; Barbier, J., Palladium and platinum-based catalysts in the catalytic

551

reduction of nitrate in water: effect of copper, silver, or gold addition. J. Catal. 2003, 220, (1),

552

182-191.

553

29. Li, H. Q.; Han, L. N.; Cooper-White, J.; Kim, I., Palladium nanoparticles decorated carbon

554

nanotubes: facile synthesis and their applications as highly efficient catalysts for the reduction of

555

4-nitrophenol. Green Chem. 2012, 14, (3), 586-591.

556

30. Kim, S. H.; Jeong, Y. E.; Ha, H.; Byun, J. Y.; Kim, Y. D., Ultra-small platinum and gold

557

nanoparticles by arc plasma deposition. Appl. Surf. Sci. 2014, 297, 52-58.

558

31. Lee, K. Y.; Byeon, H. S.; Yang, J. K.; Cheong, G. W.; Han, S. W., Photosynthesis of

559

palladium nanoparticles at the water/oil interface. Bull. Korean Chem. Soc. 2007, 28, (5), 880-

560

882.

561

32. Qin, H. M.; Qian, X. S.; Meng, T.; Lin, Y.; Ma, Z., Pt/MOx/SiO2, Pt/MOx/TiO2, and

562

Pt/MOx/Al2O3 catalysts for CO oxidation. Catalysts 2015, 5, (2), 606-633.

563

33. Brust, M.; Walker, M.; Bethell, D.; Schiffrin, D. J.; Whyman, R., Synthesis of thiol-

564

derivatized gold nanoparticles in a 2-phase liquid-liquid system. J. Chem. Soc. Chem. Commun.

565

1994, (7), 801-802.

566

34. Pol, V. G.; Srivastava, D. N.; Palchik, O.; Palchik, V.; Slifkin, M. A.; Weiss, A. M.;

567

Gedanken, A., Sonochemical deposition of silver nanoparticles on silica spheres. Langmuir

568

2002, 18, (8), 3352-3357.

569

35. Sanchez-Polo, M.; Daiem, M. M. A.; Ocampo-Perez, R.; Rivera-Utrilla, J.; Mota, A. J.,

570

Comparative study of the photodegradation of bisphenol A by HO•, SO4•- and CO3•-/HCO3•

571

radicals in aqueous phase. Sci. Total Environ. 2013, 463, 423-431.

572

36. Liang, C. J.; Huang, C. F.; Mohanty, N.; Kurakalva, R. M., A rapid spectrophotometric

573

determination of persulfate anion in ISCO. Chemosphere 2008, 73, (9), 1540-1543.

574

37. Sun, L. Z.; Bolton, J. R., Determination of the quantum yield for the photochemical

575

generation of hydroxyl radicals in TiO2 suspensions. J. Phys. Chem. 1996, 100, (10), 4127-4134.

576

38. Thompson, R. C., Catalytic decomposition of peroxymonosulfate in aqueous perchloric acid

577

by the dual catalysts Ag+ and S2O82- and Co2+. Inorg. Chem. 1981, 20, (4), 1005-1010.


Page 24 of 34

Page 25 of 34


578

39. Widmann, D.; Liu, Y.; Schuth, F.; Behm, R. J., Support effects in the Au-catalyzed CO

579

oxidation - Correlation between activity, oxygen storage capacity, and support reducibility. J.

580

Catal. 2010, 276, (2), 292-305.

581

40. Zhang, Z. Y.; Shao, C. L.; Zou, P.; Zhang, P.; Zhang, M. Y.; Mu, J. B.; Guo, Z. C.; Li, X.

582

H.; Wang, C. H.; Liu, Y. C., In situ assembly of well-dispersed gold nanoparticles on electrospun

583

silica nanotubes for catalytic reduction of 4-nitrophenol. Chem. Commun. 2011, 47, (13), 3906-

584

3908.

585

41. Elias, H.; Gotz, U.; Wannowius, K. J., Kinetics and mechanism of the oxidation of

586

sulfur(IV) by peroxomonosulfuric acid anion. Atmos. Environ. 1994, 28, (3), 439-448.

587

42. Evans, D. F.; Upton, M. W., Studies on singlet oxygen in aqueous-solution. 3. The

588

decomposition of peroxy-acids. J. Chem. Soc. Dalton Trans. 1985, (6), 1151-1153.

589

43. Serjeant, E. P.; Dempsey, B., Ionisation Constants of Organic Acids in 435 Aqueous

590

Solution. Pergamon: Oxford, U.K., 1979.

591

44. Cessna, A. J.; Grover, R., Spectrophotometric determination of dissociation constants of

592

selected acidic herbicides. J. Agric. Food Chem. 1978, 26, (1), 289-292.

593

45. Furman, O. S.; Teel, A. L.; Watts, R. J., Mechanism of base activation of persulfate.

594

Environ. Sci. Technol. 2010, 44, (16), 6423-6428.

595

46. Kwon, Y.; Lai, S. C. S.; Rodriguez, P.; Koper, M. T. M., Electrocatalytic oxidation of

596

alcohols on gold in alkaline media: Base or gold catalysis? J. Amer. Chem. Soc. 2011, 133, (18),

597

6914-6917.

598

47. Quaino, P.; Luque, N. B.; Nazmutdinov, R.; Santos, E.; Schmickler, W., Why is gold such a

599

good catalyst for oxygen reduction in alkaline media? Angew. Chem. Int. Ed. 2012, 51, (52),

600

12997-13000.

601

48. Xu, L.; Yuan, R. X.; Guo, Y. G.; Xiao, D. X.; Cao, Y.; Wang, Z. H.; Liu, J. S., Sulfate

602

radical-induced degradation of 2,4,6-trichlorophenol: A de novo formation of chlorinated

603

compounds. Chem. Eng. J. 2013, 217, 169-173.

604

49. Chan, K. H.; Chu, W., Degradation of atrazine by cobalt-mediated activation of

605

peroxymonosulfate: Different cobalt counteranions in homogenous process and cobalt oxide

606

catalysts in photolytic heterogeneous process. Wat. Res. 2009, 43, (9), 2513-2521.



607

50. Ji, Y. F.; Dong, C. X.; Kong, D. A.; Lu, J. H., New insights into atrazine degradation by

608

cobalt catalyzed peroxymonosulfate oxidation: Kinetics, reaction products and transformation

609

mechanisms. J. Hazard. Mater. 2015, 285, 491-500.

610

51. Kim, S.; Choi, W., Visible-light-induced photocatalytic degradation of 4-chlorophenol and

611

phenolic compounds in aqueous suspension of pure titania: demonstrating the existence of a

612

surface-complex-mediated path. J. Phys. Chem. B. 2005, 109, (11), 5143-9.

613

52. Lee, H.; Yoo, H. Y.; Choi, J.; Nam, I. H.; Lee, S.; Lee, S.; Kim, J. H.; Lee, C.; Lee, J.,

614

Oxidizing capacity of periodate activated with iron-based bimetallic nanoparticles. Environ. Sci.

615

Technol. 2014, 48, (14), 8086-8093.

616

53. Wang, Y. X.; Sun, H. Q.; Ang, H. M.; Tade, M. O.; Wang, S. B., 3D-hierarchically

617

structured MnO2 for catalytic oxidation of phenol solutions by activation of peroxymonosulfate:

618

Structure dependence and mechanism. Appl. Catal. B Environ. 2015, 164, 159-167.

619

54. Fang, G. D.; Gao, J.; Dionysiou, D. D.; Liu, C.; Zhou, D. M., Activation of persulfate by

620

quinones: Free radical reactions and implication for the degradation of PCBs. Environ. Sci.

621

Technol. 2013, 47, (9), 4605-4611.

622

55. Zhou, Y.; Jiang, J.; Gao, Y.; Ma, J.; Pang, S. Y.; Li, J.; Lu, X. T.; Yuan, L. P., Activation of

623

peroxymonosulfate by benzoquinone: A novel nonradical oxidation process. Environ. Sci.

624

Technol. 2015, 49, (21), 12941-12950.

625

56. Ingram, A. J.; Walker, K. L.; Zare, R. N.; Waymouth, R. M., Catalytic role of multinuclear

626

palladium-oxygen intermediates in aerobic oxidation followed by hydrogen peroxide

627

disproportionation. J. Amer. Chem. Soc. 2015, 137, (42), 13632-13646.

628

57. Ball, D. L.; Edwards, J. O., The catalysis of the decomposition of Caro's acid. J. Amer.

629

Chem. Soc. 1958, 62, 343-345.

630

58. Lee, J. S.; Choi, W. Y., Photocatalytic reactivity of surface platinized TiO2: Substrate

631

specificity and the effect of Pt oxidation state. J. Phys. Chem. B. 2005, 109, (15), 7399-7406.


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Page 27 of 34


FIGURE 1. Representative TEM images of (a) palladium, (b) platinum, (c) gold, and (d) silver particles loaded on alumina supports.



0.30

Page 28 of 34

100 4-CP Conc. Pd-Al2 O3 /PMS

1.0 80

0.20

0.15

0.10

0.05

60 0.6 40

0.4

20

0.2

Pd-TiO2 /PMS Chloride Ion Conc. ( µM)

0.8 4-CP Conc. (C/C0 )

PMS or Sulf ate Ion Conc. (mM)

0.25

Chloride ion Conc. Pd-Al2 O3 /PMS Pd-TiO2 /PMS PMS Conc. Pd-Al2 O3 /PMS Pd-TiO2 /PMS Sulf ate ion Conc. Pd-Al2 O3 /PMS Pd-TiO2 /PMS

0.00

0.0

0 0

10

20

30

40

50

60

Time (min)

FIGURE 2. Degradation of 4-CP, reduction of PMS, formation of sulfate ions, and evolution of chloride ions (as a result of 4-CP dechlorination) by Pd-Al2O3 and Pd-TiO2 combined with PMS ([Pd-Al2O3]0 = [Pd-TiO2]0 = 0.25 g/L; [PMS]0 = 0.25 mM; [4-CP]0 = 0.1 mM; [bicarbonate]0 = 1 mM; pHi = 7.0). Concentration of sulfate ions was calculated by subtracting 0.25 mM from the measured sulfate concentration (sulfate ions were initially present due to addition of 0.25 mM PMS (i.e., KHSO5•0.5KHSO4•0.5K2SO4)).


Page 29 of 34


20 4-CP Conc. Pd-Al2 O3

8

15

6 10 4 5 2

0

PMS Reduction Rate Rate ( µM/min)

Initial Rate of 4-CP Degradation ( µM/min)

10

Pt-Al2 O3 Au-Al2 O3 PMS Conc. Pd-Al2 O3 Pt-Al2 O3 Au-Al2 O3

0 0.125 g/L

0.25 g/L

0.375 g/L

0.5 g/L

FIGURE 3. Initial rate of 4-CP degradation in the presence of PMS by Pd-, Pt-, and Au-Al2O3 systems as a function of the activator loading ([NM-Al2O3]0 = 0.125, 0.25, 0.375, and 0.5 g/L; [PMS]0 = 0.25 mM; [4-CP]0 = 0.1 mM; [bicarbonate]0 = 1 mM; pHi = 7.0).



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14 4-CP Conc. Pd-Al2 O3

12 8 10 6

8 6

4 4 2

Pt-Al2 O3 PMS Reduction Rate ( µM/min)

Initial Rate of 4-CP Degradation ( µM/min)

10

Au-Al2 O3 Ag-Al2 O3 PMS Conc. Pd-Al2 O3 Pt-Al2 O3 Au-Al2 O3 Ag-Al2 O3

2 0

0 pH 3

pH 6

pH 7 (b uffered)

pH 9

pH 11

FIGURE 4. Effect of initial pH on the kinetic rate of 4-CP degradation and PMS reduction in the aqueous suspensions of Pd-, Pt-, Au-, and Ag-Al2O3 ([NM-Al2O3]0 = 0.25 g/L; [PMS]0 = 0.25 mM; [4-CP]0 = 0.1 mM; pHi = 3, 6, 7 (buffered using 1 mM bicarbonate), 9, and 11).


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Pd-Al2O3

100

Pt-Al2O3 Au-Al2 O3 Ag-Al2 O3

Degradation Eff iciency ( %)

80

Pd-TiO2 Pt-TiO2

60

40

20

0 BPA

PH

BZA

4-NP

CBZ

4-CP

2,4,6-TCP

FIGURE 5. Oxidative degradation of various organic compounds by PMS activated with noble metals loaded on alumina and titania ([NM-Al2O3]0 = [NM-TiO2)]0 = 0.25 g/L; [PMS]0 = 0.25 mM; [bisphenol A (BPA)]0 = [phenol (PH)]0 =[benzoic acid (BZA)]0 = [4-nitrophenol (4-NP)]0 = [carbamazepine (CBZ)]0 = [4-chlorophenol (4-CP)]0 = [2,4,6-trichlorophenol (2,4,6-TCP)]0 = 0.1 mM; [bicarbonate]0 = 1 mM; pHi = 7.0).



(a)

Pd-Al2O3 /PMS

1.0

w /o MeOH w / MeOH Au-Al2O3 /PMS

150

0.8

Co 2+/PMS HCHO Conc . (µ M)

4-CP Conc. (C/C0)

Page 32 of 34

0.6

w /o MeOH w / MeOH 2+ Co /PMS w /o MeOH w / MeOH

100

50 Pd-Al 2O3/PMS

0.4

Au-Al 2O3/PMS

0 0

5

10

15

T ime (min)

0.2

0.0 0

10

20

30

40

50

60

2+

(b) 6

3

5 4

2

3 2

1

1 0

Rate of Substrate Degradation by MWCNT/PMS and Pd-Al2O3/PMS ( µM/min)

Rate of Substrate Degradation by Co2+/PMS ( µM/min)

Time (min) 7 Co /PMS MWCNT/PMS Pd-Al2O3 /PMS

0 BZA CBZ

BPA 4-NP

PH

4-CP 2,4,6-TCP

FIGURE 6. (a) Effect of methanol on the rate of 4-CP degradation by the Pd-Al2O3/PMS, AuAl2O3/PMS, and Co2+/PMS systems ([Pd-Al2O3]0 = [Au-Al2O3]0 = 0.25 g/L; [Co2+]0 = 25 µM; [PMS]0 = 0.25 mM; [4-CP]0 = 0.1 mM; [methanol]0 = 0.25 M; [bicarbonate]0 = 1 mM; pHi = 7.0). Inset: oxidative conversion of methanol to formaldehyde ([Pd-Al2O3]0 = [Au-Al2O3]0 = 0.25 g/L; [Co2+]0 = 25 µM; [PMS]0 = 1 mM; [methanol]0 = 1 mM; [bicarbonate]0 = 1 mM; pHi = 7.0) (b) Comparison of PMS activators (i.e., Co2+, MWCNT, and Pd-Al2O3) in terms of substrate specificity ([Pd-Al2O3]0 = [MWCNT]0 = 0.25 g/L; [Co2+]0 = 0.25 mM; [PMS]0 = 0.25 mM; [target substrate]0 = 0.1 mM; pHi = 3.85). 32 Environment ACS Paragon Plus

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1st

2nd

3rd

4th

5th

Pd-Al2 O3

1.0

Pt-Al2 O3 Pd-TiO2

4-CP Conc. (C/C0 )

0.8

Pt-TiO2

0.6

0.4

0.2

0.0 0

50

100

150

200

250

300

Time (min)

FIGURE 7. Repeated degradation of 4-CP by Pd and Pt deposited on alumina and titania supports ([Pd-Al2O3]0 = [Pt-Al2O3]0 = [Pd-TiO2]0 = [Pt-TiO2]0 = 0.25 g/L; [PMS]0 = 0.25 mM; [4-CP]0 = 0.1 mM; [bicarbonate]0 = 1 mM; pHi = 7.0).



Table of Contents Figure:


Page 34 of 34

Activation of Peroxymonosulfate by Surface-Loaded Noble Metal

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