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Accelerated catalytic Fenton reaction with traces of iron – An Fe-Pd-multicatalysis approach Anett Georgi, Miriam Velasco Polo, Klara Crincoli, Katrin Mackenzie, and Frank-Dieter Kopinke Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.6b01049 • Publication Date (Web): 11 May 2016 Downloaded from http://pubs.acs.org on May 26, 2016

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Environmental Science & Technology

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Accelerated catalytic Fenton reaction with traces of iron – An Fe-Pd-multicatalysis

2

approach

3

Anett Georgi*, Miriam Velasco Polo, Klara Crincoli, Katrin Mackenzie and Frank-Dieter Kopinke

4

Helmholtz Centre for Environmental Research - UFZ, Department of Environmental

5

Engineering, Permoserstr. 15, D-04318 Leipzig, Germany

6 7

Abstract

8

An accelerated catalytic Fenton (ACF) reaction was developed based upon a multicatalysis

9

approach, facilitating efficient contaminant oxidation at trace levels of dissolved iron. Beside

10

the FeII/H2O2 catalyst/oxidant pair for production of OH-radicals, the ACF system contains

11

Pd/H2 as catalyst/reductant pair for fast reduction of FeIII back to FeII which accelerates the

12

Fenton cycle and leads to faster contaminant degradation. By this means, the concentration of

13

the dissolved iron catalyst can be reduced to trace levels (1 mg L-1) below common discharge

14

limits, thus eliminating the need for iron sludge removal, which is one of the major drawbacks

15

of conventional Fenton processes. ACF provides fast degradation of the model contaminant

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methyl tert-butyl ether (MTBE, C0 = 0.17 mM) with a half-life of 11 min with 1 mg L-1 dissolved

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iron, 500 mg L-1 H2O2, 5 mg L-1 Pd (as suspended Pd/Al2O3 catalyst) and 0.1 MPa H2, pH = 3. The

18

effects of pH, H2 partial pressure and H2O2 concentration on MTBE degradation rates were

19

studied. Results on kinetic deuterium isotope effect and quenching studies are in conformity

20

with OH-radicals as main oxidant. The heterogeneous Pd/Al2O3 catalyst was reused within 6

21

cycles without significant loss in activity.

22

TOC art

Heterogeneous catalyst/reductant

Homogeneous

+ catalyst/oxidant

= Multicatalysis for ACF

23 24

1. Introduction

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Advanced oxidation processes (AOPs) are an indispensable tool for removing toxic, hardly

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biodegradable and recalcitrant organic contaminants from water. The Fenton reaction (Eq. 1) 1 ACS Paragon Plus Environment


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allows the OH-radical, one of the most powerful oxidants (E◦ = 2.73 V, pH = 3), to be produced

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at ambient conditions from iron salts and H2O2. These reagents are inexpensive, relatively easy

29

to store and handle, and environmentally benign.

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However, the homogeneous Fenton reaction with dissolved iron as catalyst also has various

31

disadvantages, mainly related to iron speciation. The formation of Fe(III) oxyhydroxides during

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the reaction must be prevented by means of acidification of the incoming water (optimal pH

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around 3)1. After the final neutralization of the treated water, iron sludge is inevitably

34

produced and must be disposed of. Heterogeneous Fenton-like catalysts offer improved

35

properties for catalyst recycling. However, at the same time they often suffer from disadvan-

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tages such as relatively low catalytic activity and utilization efficiency of H2O2, as reported for

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iron oxides2-4, or insufficient stability and leaching of metal ions.5 Furthermore, catalysts with

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iron on microporous supports such as zeolites suffer steric limitations for the access of large

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pollutant molecules.6

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In order to at least reduce the amount of iron sludge, the so-called catalytic Fenton or Fenton-

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like reaction is used, whereby FeII or FeIII salts are applied in sub-stoichiometric, i.e. catalytic

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amounts (typically in the range of 10–100 mg L-1 1,7) related to the amount of H2O2. After the

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initial fast consumption of the FeII added, the slow reduction of FeIII back to FeII by reaction

44

with H2O2 (Eq. 2) or HO2●/O2●− (Eq. 3) is the rate-limiting step in these systems. Despite the

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comparably high rate constant of the latter reaction, its contribution is limited by the low

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stationary concentration of the transient radicals. Further important reactions of the complex

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Fenton chemistry together with their second-order rate constants at pH = 3 are listed below:8

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FeII + H2O2 → FeIII + ●OH + OH−

k = 63 M-1 s-1

Eq. 1

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FeIII + H2O2 → FeII + HO2● + H+

k = 2.0 × 10−3

Eq. 2

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FeIII + HO2●/O2●− → FeII + O2 + (1 or 0) H+

k = 7.8 × 105

Eq. 3

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H2O2 + ●OH → HO2● + H2O

k = 3.3 × 107

Eq. 4

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FeII + ●OH → FeIII + OH−

k = 3.2 × 108

Eq. 5

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FeII + HO2●/O2●− + (1 or 2) H+ → FeIII + H2O2

k = 1.3 × 106

Eq. 6

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For simplification, the terms FeIII and FeII are used in the following sections for all iron species

55

of these oxidation states present in aqueous solution, irrespective of their ligands. It has been

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suggested that the reaction between FeIII and H2O2 described by the net reaction in Eq. 2

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proceeds via an initial complex formation (Fe3+ + H2O2 ⇋ Fe(HO2)2+ + H+) and subsequent

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reductive dissociation of the FeIII-peroxide complex (Fe(HO2)2+ → Fe2+ + HO2●,1 as described in

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further detail in the Supporting Information (SI, part S2).

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The reduction of FeIII can be enhanced by applying UV light (photo-Fenton9,10), electrical

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potential at suitable electrodes (electro-Fenton11) or additional chemical reductants or redox

62

mediators (for references see below). However, photo-Fenton processes require the

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availability of intense solar irradiation or energy-consuming UV lamp irradiation. For electro-

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Fenton processes, achieving high space-time yields is a challenge which limits their applicability

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for large water volumes. The addition of suitable reducing agents or redox mediators to

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enhance FeIII/FeII recycling thus offers an interesting alternative. Optimal reagents should

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significantly increase the rate of FeIII reduction without leaving undesired residues in the

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treated water and, in addition, have at most a minor impact as parasitic consumers of the

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reactive oxidant species produced.

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In previous studies, humic acid12 and quinones13 have been suggested as redox mediators

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which create an intermediate electron-shuttling cycle between FeIII as final electron acceptor

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and H2O2 as electron donor. In addition, hydroxylamine14 and ascorbic acid15 have been applied

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in excess to FeIII in order to provide a stronger reducing agent in addition to H2O2. All these

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compounds, however, are themselves attacked by ●OH and thus consumed during the

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reaction. Even more important is that their addition may lead to residual degradation

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products, such as increased levels of dissolved organic carbon (DOC) in the case of incomplete

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mineralization of organic mediators, or nitrate in the case of hydroxylamines.14

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H2 in comparison is a clean reductant, leaving only water as product. In addition, its reactivity

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towards ●OH is relatively low (k•OH,H2 = (3.4–6.0) x 107 M-1 s-1).16 Dissolved H2 is, however,

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inactive for reduction of dissolved FeIII to FeII at ambient conditions. Taking into account the

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dominant iron species at pH = 3 ([Fe(H2O)5(OH)]2+)17, this reaction (Eq. 7) is thermodynamically

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favorable (∆ < 0), but obviously kinetically limited due to high activation energy barriers. 1 [Fe H O) OH)] + H → [Fe H O) ] 2

∆ = −87 kJ mol"#

Eq. 7

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Typically, noble metal catalysts such as Pd are used in order to activate H2 and the combined

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application of H2 and Pd catalysts proved successful for water treatment by reductive

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transformation of contaminants even at large scale.18 Thus, we hypothesize that activated

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hydrogen formed on noble metal catalysts such as Pd can be utilized to reduce dissolved FeIII to

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FeII (Eqs. 8 and 9). 2 Pd + H → 2 Pd ∙ H

Eq. 8

Pd ∙ H + Fe''' → Pd + H + Fe''

Eq. 9

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Pd has also been shown to catalyze the formation of H2O2 from O2 and H2 (Eq. 11).19 Yuan et al.

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have designed a process whereby: I) O2 and H2 are produced by H2O electrolysis, II) H2O2 is

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produced from these gases on Pd catalysts, and III) dissolved iron is responsible for production 3 ACS Paragon Plus Environment


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of ●OH from H2O2.20,21 For example, Rhodamine B oxidation by this process was performed

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with 50 mg L-1 dissolved iron.20 The involvement of the Pd catalyst in the reduction of dissolved

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FeIII was not considered in these papers. Although Pd was suggested to catalyze, to some

94

extent, the formation of ●OH from H2O2 (Eq. 11), its efficiency for this reaction is low.20,22 Since

95

Pd also catalyzes the decomposition (Eq. 12) and hydrogenation of H2O2 (Eq. 13) as well as H2

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combustion (Eq. 14), the selectivity for H2O2 formation is limited. Consequently, accumulated

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H2O2 concentrations obtained in the studies cited above were rather low (< 20 mg L-1 in 40

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min).20,21 In addition, activated hydrogen can also act as a consumer of ●OH, formed from H2O2

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e.g. in a Fenton reaction in the solution phase (Eq. 15). ()

H + O *+ H O

Eq. 10

()

H O *+ 2 ●OH

Eq. 11

() 1 H O *+ H O + O 2

Eq. 12

()

H O + H *+ 2 H O

Eq. 13

1 () H + O *+ H O 2

Eq. 14

Pd ∙ H + ●OH → Pd + H O

Eq. 15

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In another setup applied for in situ formation of H2O2 and Fenton oxidation, the Pd catalyst

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was supported on magnetite.23 Considerable amounts of dissolved iron were produced under

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the applied acidic conditions (pH ≤ 3, CFe,dissolved ≥ 10 mg L-1 within 60 min corresponding to ≈

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95% phenol degradation). This was ascribed to reductive dissolution of magnetite by the Pd-

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catalyzed formation of activated hydrogen.23 Electrochemical production of Fe2+ by an iron

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cathode allowed better control of dissolved iron concentration by adjusting the current

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applied to the cathode resulting in 91% MTBE turnover (k’ = 0.04 min-1, C0,MTBE = 20 mg L-1)

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within 60 min with simultaneous production of 14 mg L-1 dissolved iron.24

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In contrast to the cited literature studies, the aim of the present study was to design an

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accelerated catalytic Fenton (ACF) reaction with minimal concentrations of dissolved iron, thus

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eliminating the need for sludge removal as a post-treatment step of the Fenton oxidation. In

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most countries, discharge limits for total iron are in the range of 0.5 to 2 mg L-1.25 Thus, the

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ACF system was intended to work at these low Fe concentrations. The suggested process relies

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on the acceleration of FeIII/FeII recycling by introducing H2 as reductant and supported Pd as a

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second catalyst. A pre-requisite for this concept is the reduction of dissolved FeIII by activated

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hydrogen at the Pd surface even in the presence of H2O2: a process which, to the best of our

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knowledge, has not previously been experimentally proven. If existent, this process should also

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be relevant for Fenton systems with in situ H2O2 formation, since they contain all required 4 ACS Paragon Plus Environment

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components (i.e. Pd, H2 and dissolved FeIII), even though catalytic reduction of FeIII to FeII at the

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Pd surface has not yet been considered in mechanistic schemes of such studies.20,21,26,27

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Methyl tert-butyl ether (MTBE) was applied as model contaminant, due to its relevance as

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groundwater contaminant together with its inertness regarding losses by processes other than

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radical-driven oxidation (i.e. low volatilization and sorption tendency, inertness towards

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reduction). MTBE has a moderate reactivity towards ●OH (k•OH,MTBE = 1.6 x 109 M-1 s-1).16 H2 and

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H2O2 were added externally.

125 126

2. Materials and Methods 2.1. Chemicals and Materials

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MTBE, tert-butyl formate (TBF), tert-butyl alcohol (TBA), FeSO4•7H2O, Na2SO3 and acetone, all

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in high purity (≥ 98%), as well as H2O2 (30 wt-%) were obtained from Merck, Germany. MTBE-

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d12 with 99 atom-% deuterium was obtained from Sigma, Germany and TiOSO4 solution (1.9–

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2.1 wt-%) from Fluka, Germany.

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As Pd catalyst, the sieve fraction 63–125 µm of the crushed G-133D Pd on ɣ-Al2O3 egg-shell

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catalyst with 0.5 wt-% Pd from Commercia, Germany, was used. The catalyst has a BET surface

133

area of about 160 m2 g-1 and a Pd dispersion of about 20%.

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2.2 Procedure of oxidation experiments

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The reactions were carried out in 150 mL Erlenmeyer flasks equipped with Mininert® valves.

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For ACF experiments, a defined amount of catalyst (50 mg catalyst, corresponding to 5 mg L-1

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Pd) was suspended in 50 mL aqueous solution (pH = 3 adjusted with HNO3) containing the

138

required amount of FeII (18–36 µM), which was added from a freshly prepared stock solution

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of FeSO4. Where not otherwise stated, the solution was purged with H2 or H2/N2 gas mixtures

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for 10 min in order to fill the headspace of the reaction vessel with an appropriate H2 reservoir

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(3.3 mmol in case of pure H2 atmosphere) and to equilibrate it with the aqueous reaction

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suspension. After that the vessel was closed and the desired concentration of MTBE was added

143

from an aqueous stock solution. For reasons of comparison, a set of blank experiments was

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conducted analogously without addition of one of the catalysts (Pd or Fe) or with various gas

145

compositions (H2/N2 and H2/O2 mixtures or air). For experiments on the kinetic isotope effect

146

of MTBE oxidation, MTBE-d0 and -d12 were applied in equimolar amounts in one and the same

147

reaction batch. In all cases, the reaction was started by adding H2O2. In some experiments H2O2

148

was re-dosed at certain time intervals as indicated. If mentioned (only for experiments with

149

Cgas,H2 ≈ 5 vol-%, Figure 5b), H2 re-dosing was done by injecting a small volume (1 mL) of pure H2

150

into the headspace (80 mL) of the reaction vessel. The vessel was continuously shaken (210

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min-1) or stirred by a magnetic stirrer (500 rpm) throughout the whole experiment, kept at

152

ambient temperature (T = 23 ± 2°C) and either kept closed or equipped with a ‘breathing 5 ACS Paragon Plus Environment


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system’ (pressure equalization by temporary connection to a continuously purged N2 reservoir,

154

see Fig. S1 in SI). The breathing system was applied especially for analyzing the change in gas

155

phase composition during the reaction and avoiding under-pressure (due to H2 consumption).

156

By means of experiments with variation of agitation intensity it was verified that reaction rates

157

were not controlled by external mass transfer effects including H2 delivery from the headspace

158

into the solution phase. The pH of the reaction suspension was found to be constant within the

159

range of pH = 3.0 ± 0.2. At given time intervals, liquid samples for H2O2 and MTBE analysis

160

were taken. Samples (100 µL) for MTBE analysis by Headspace-GC-MS were filled into 10 mL

161

vials containing 2 mL of 5 mM Na2SO3 in deionized water for quenching of residual H2O2.

162

Recycling experiment: After each run, the bottles containing catalyst and reaction solution

163

were treated with ultrasound for 30 minutes then centrifuged at 3000 rpm for 60 min, after

164

which the liquid phase (~50 mL) was largely removed and replaced by the same volume of

165

deionized water, whereby the solid catalyst was transferred again into the reaction vessel. The

166

next reaction cycle was started as described above.

167

Stoichiometric Fenton experiment: This experiment was conducted in order to determine the

168

hydrogen kinetic isotope effect for MTBE degradation by ●OH. A solution containing 15 mM FeII

169

and 0.2 mM each of MTBE-d0 and -d12 was adjusted with HNO3 to pH = 3. Under vigorous

170

stirring, 15 mM H2O2 were added. After 4 min the reaction was stopped by adding 30 mM

171

Na2SO3; the residual concentrations of MTBE-d0 and -d12 were determined by means of

172

Headspace-GC-MS.

173

Data shown in the figures are mean values of at least three experiments. The error represents

174

the mean deviation of the single values from the mean value.

175

2.3. Analytical methods

176

MTBE-d0 and -d12, TBF, TBA and acetone were analyzed by means of Headspace-GC-MS using a

177

QP 2010 GC-MS device equipped with an AOC-5000 autosampler (Shimadzu Corp.) and a

178

Zebron ZB-5MSi capillary column (30 m x 0.25 mm x 0.5 µm, Phenomenex, Germany).

179

Quantification was performed in SIM mode using the following characteristic ions: m/z = 73 for

180

MTBE-d0, 82 for MTBE-d12, 58 for acetone and 59 for TBA and TBF.

181

H2O2 concentrations were measured photometrically (λ = 405 nm) after adding titanyl sulfate

182

solution. FeII was determined by means of the phenanthroline method28 adapted to analysis of

183

low concentrations (4.5–36 µM) by using 100 QS/50 mm quartz cuvettes. For determining total

184

dissolved iron concentrations, all FeIII present was reduced by addition of ascorbic acid before

185

phenanthroline was added.


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Gas analyses for determination of H2, N2 and O2 were conducted with a GC-TCD system

187

(Agilent 6850) equipped with an HP plot column (30 m x 0.32 mm x 12 µm, Agilent), oven

188

temperature 30°C.

189 190

3. Results and discussion 3.1. FeIII reduction by Pd/H2

191

Initially, the rate of FeIII reduction by the Pd/alumina catalyst in the presence of H2 at pH = 3

192

was studied. With 5 mg L-1 Pd, conversion of FeIII (2 mg L-1 or 36 µM) into FeII was fast (>80%

193

within 1 min, Fig. 1). The recovery of total dissolved iron was within (100 ± 5)% before and

194

after the reduction step, indicating that adsorption of dissolved FeII and FeIII to the solid

195

catalyst was negligible under the applied conditions. At a 10-fold lower Pd concentration (0.5

196

mg L-1), the half-life (t0.5) of FeIII reduction was about 6 min, which implies a specific catalytic

197

activity APd of 330 L g-1 min-1 for FeIII reduction by the Pd/alumina catalyst according to APd =

198

1/(t0.5 × CPd).29 This activity is in the upper range reported for various H2-based reduction

199

reactions catalyzed by supported Pd catalysts (summarized by Chaplin et al.18) including

200

inorganic ions and halogenated organic compounds as substrates. Possibly, the observed rates

201

are already partially mass transfer limited such that the ‘true’ reaction rates are even higher.

202

For comparison, the highly efficient hydrodechlorination of various chlorinated ethenes by the

203

same Pd catalyst (pH = 3) as used in this study runs with Pd activities in the range of 200 to

204

1200 L g-1 min-1.30 Obviously, there is a possibility that activated hydrogen provided by the Pd

205

catalyst can contribute to FeIII reduction in the Fenton system.

206

207 208

Figure 1: FeIII reduction by Pd/H2 (CFeIII,0 = 36 µM, pH = 3, CPd = 0.5 or 5 mg L-1, respectively, as

209

Pd/Al2O3 (0.5 wt-% Pd), Vwater = 50 mL, Vgas = 80 mL (100 vol-% H2)).



210

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3.2. Proof of principle

211

Based on the observation that FeIII can be rapidly reduced by H2 in the presence of a

212

heterogeneous Pd catalyst, we conducted an experiment for degradation of MTBE (CMTBE,0 =

213

0.17 mM) under ACF conditions, i.e. in addition to H2O2 (CH2O2,0 = 15 mM) and FeII (CFeII,0 = 1 mg

214

L-1), the reaction suspension contained 5 mg L-1 Pd (as Pd/Al2O3) and H2 was provided as

215

reductant (pH2 = 0.1 MPa). Fig. 2a shows that MTBE was degraded by (97 ± 1)% within 60 min,

216

following pseudo-first-order kinetics with k’ = (0.061 ± 0.005) min-1 and a half-life of about 11

217

min in the ACF system. Control batch experiments were carried out in order to clarify the role

218

of H2 and Pd in the system. Fig. 2a shows that in a conventional catalytic Fenton (CF) system

219

under identical conditions, i.e. CFeII,0 = 1 mg L-1, CH2O2,0 = 15 mM and CMTBE,0 = 0.17 mM, MTBE

220

degradation is extremely slow (≤10% turnover within 90 min). This result is to be expected

221

based on previous studies where the dosage of Fe salts to achieve high MTBE degradation

222

extents within 1-2 h by homogeneous Fenton systems was in the range of 10-260 mg L-1,31-36

223

with the lower limit being for the case of a very low MTBE concentration (0.02 mM).31,32 The

224

addition of only H2 to the CF system has no effect on MTBE degradation rates. The addition of

225

only the Pd catalyst to the CF system leads to a slow but significant MTBE degradation with k’ =

226

(0.0048 ± 0.0005) min-1. The Pd catalyst without dissolved iron showed no significant activity

227

for MTBE oxidation with H2O2, which confirms that OH-radical formation from H2O2 by Pd

228

alone (Eq. 11) is of only minor importance. Thus, there already seems to be a small promoting

229

effect of Pd and Fe even in the absence of H2. Elucidation of the underlying mechanisms of this

230

moderate effect is out of the scope of the present study; some discussion is provided in the SI

231

(part S2). In contrast to the above, a tremendous rate enhancement in the degradation of

232

MTBE is observed in the complete ACF system containing Pd and dissolved FeII as catalysts

233

together with H2 as reductant and H2O2 as oxidant (k’ = 0.061 vs. 0.0048 min-1). Overall, these

234

results emphasize the necessity of the Pd catalyst for the activation of H2 by dissociative

235

adsorption (Eq. 8), providing activated hydrogen as a new reducing agent for FeIII (Eq. 9). The

236

reduction of FeIII by activated hydrogen replaces the slow step of FeIII reduction by H2O2 and

237

thus strongly accelerates the Fe-redox cycle. In accordance with this hypothesis, H2O2

238

consumption (Fig. 2b) is also fastest in the ACF system.

239

It is remarkable that MTBE degradation follows an apparent first-order kinetics down to very

240

low residual concentrations (C/C0 ≤ 0.004), although H2O2 is consumed at a comparable rate

241

(k’MTBE = 0.061 min-1 vs. k’H2O2 = 0.048 min-1). This means that H2O2 is not involved in the rate-

242

limiting step of the ACF oxidation. This conclusion is in conformity with the hypothesis that the

243

hydrogen-driven reduction of FeIII is the rate-limiting step, even under ACF conditions.



ACF CF+Pd

6.5

CF Pd

110

- ln (C/C0) MTBE

5.5

90

4.5

70

3.5

y = 0.061x

50

2.5

30

1.5

y = 0.0048x

0.5

y=0x

-0.5

-10 0

20

40

A

60

80

100

t in min ACF CF + Pd Pd

2.2 1.8

-ln (C/C0) H2O2

10

MTBE conversion in %

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CF CF + H2 Pd + H2 y = 0.048x R² = 0.9996

1.4

y = 0.014x R² = 0.996 y = 0.0094x R² = 0.995 y = 0.0072x R² = 0.998

1.0 0.6 0.2

y = 0x

-0.2 0

B

20

40

60

80

100

t in min

244 245 246 247 248 249 250 251 252

Figure 2 a): MTBE degradation and 2 b): H2O2 consumption under accelerated Fenton (ACF) conditions in comparison to conventional catalytic Fenton (CF) conditions with/without amendment by Pd or H2 individually and in Fe-free Pd-catalyzed reactions. All reactions were conducted in presence of MTBE and H2O2 under the following conditions: CH2O2,0 = 15 mM, CMTBE,0 = 0.17 mM, pH = 3, Vwater = 50 mL, Vgas = 80 mL. Variable components (listed behind symbols and graph notation) were gas composition, dissolved iron (if added: CFe = 18 µM), and Pd/Al2O3 (if added: CPd = 5 mg L-1). (■) CF: FeII, air atmosphere; (▲) ACF: FeII, Pd, H2 atmosphere; (◆) CF + Pd: FeII, Pd, air atmosphere; (×) Pd: Pd, air atmosphere; (ο) CF + H2: FeII, H2 atmosphere; (+) Pd + H2: Pd, H2 atmosphere.

253

In the CF system, H2O2 decomposition is slow due to the low concentration of the Fe catalyst

254

and the bottleneck of its recycling to FeII. H2O2 decomposition in the ACF system can occur via

255

several pathways: I) the regular Fenton chemistry (i.e. via Eqs. 1, 2, 4) which is enhanced by the

256

accelerated Fe(II/III) cycle; II) catalytic dissociation of H2O2 at the Pd surface (Eq. 12), and III) 9 ACS Paragon Plus Environment


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hydrogenation by activated hydrogen at the Pd surface (Eq. 13). In order to estimate the

258

contribution of the various pathways, we conducted individual experiments with the ACF

259

components. The addition of H2 to the CF system had no effect on H2O2 consumption. When

260

Pd alone is present, or only Pd is added to the CF system, H2O2 consumption is already

261

significant due to the Pd-catalyzed H2O2 decomposition (Eq. 12). The presence of H2 further

262

accelerates H2O2 consumption via its hydrogenation over Pd (Eq. 13). This is in accordance with

263

the finding that catalytic H2O2 decomposition (Eq. 12) and hydrogenation (Eq. 13) are equally

264

important in Pd/H2/H2O2 systems at acidic pH.19 Nevertheless, the addition of 1 mg L-1 of

265

dissolved iron to the Pd/H2/H2O2 system, i.e. applying the complete ACF system, still strongly

266

accelerates H2O2 decomposition indicating that the largest part of H2O2 decomposition in the

267

ACF system is caused by the accelerated Fe redox cycle driven by reduction of FeIII at the Pd

268

surface. It is remarkable that the FeIII reduction rate can successfully compete with the H2O2

269

reduction rate, although the H2O2 concentration is about three orders of magnitude higher

270

(e.g. 15 mM vs. 18 µM, in Fig. 2).

271

In any Fenton system, the overall efficiency of H2O2 utilization for contaminant degradation is

272

the product of the efficiencies of ●OH production from H2O2 (E1) and the efficiency of ●OH

273

consumption by reaction with the target contaminant (E2) as described by Eqs. 16 and 178,37 for

274

MTBE. , -MTBE ) , -H2O2 )

=

moles of ●OH produced moles of H2O2 consumed

×

moles of MTBE consumed moles of ●OH produced

= 31 × 32

η = 3# × 3 × 100%

Eq. 16

Eq. 17

275

As mentioned above, H2O2 consumption in the ACF system is composed of the productive

276

pathway leading to ●OH (Eq. 1) and non-productive parallel reactions including those at the Pd

277

surface (Eqs. 12 and 13). The contribution of the latter can be estimated based on the H2O2

278

decomposition rate constant determined in the presence of Pd and H2 without dissolved iron

279

(reaction Pd + H2 in Figure 2b, k’Pd+H2 = 0.014 min-1). Thus, based on the rate laws for parallel

280

reactions, the contribution of H2O2 consumption at the Pd catalyst (Eqs. 12 and 13) to its

281

overall turnover (X) in the ACF system accounts for only 30% (XPd+H2/Xtotal,ACF = k’Pd+H2/k’total,ACF ≈

282

0.3). This non-productive consumption reduces the available amount for ●OH production,

283

leading to an estimated reduction in E1 by a factor of 0.7. It is worthy to note that in all

284

catalytic Fenton reactions FeIII reduction consumes H2O2 which is not available for ●OH

285

production (unless organic reaction intermediates significantly contribute to FeIII reduction).

286

Depending on the reaction conditions and the assumed pathway of the catalytic Fenton cycle,

287

the theoretical efficiency of ●OH production under ideal conditions was calculated to be

288

between 0.58 to 0.67.37 The loss of H2O2 as reductant of FeIII in the ‘conventional’ catalytic

289

Fenton system is at least as high as its parasitic consumption in the ACF system. The factor E2

290

in Eq. 16 is the result of the competition among various consumers for ●OH, including the 10 ACS Paragon Plus Environment

Page 11 of 23


291

target compound (MTBE), its degradation intermediates, H2O2, the Pd/Al2O3 catalyst (with

292

activated H) and possibly other transient species. In a well-defined homogeneous system E2

293

can be estimated from the second-order rate constants and concentrations of all identified

294

●

295

unknown and thus this estimation is not straightforward. Nevertheless, it is obvious that the

296

concentration of H2O2 needs to be optimized as it is one of the relevant consumers of ●OH.

297

Thus, the overall efficiency of H2O2 utilization for MTBE degradation (η, Eq. 17) will be

298

discussed in section 3.4.

299

Even though quantification of steady-state FeII concentrations in the ACF system would be

300

desirable, it cannot be realized with standard laboratory equipment. Any phase separation

301

(removal of solid Pd catalyst) or FeII/III complexation (e.g. with phenanthroline for photo-

302

metry28,38) affects the steady-state FeII concentration, since it is controlled by a number of fast

303

reactions. Nevertheless, we were able to show on a semi-quantitative level that significant FeII

304

concentrations exist in the aqueous phase over the time course of an ACF experiment,

305

whereas in the absence of Pd and H2 (i.e. in a normal catalytic Fenton reaction with the same

306

initial FeII and H2O2 concentrations), FeII concentration dropped below the detection limit

307

within 5 min (SI, part S3). This result confirms again that the reduction of FeIII to FeII by

308

activated hydrogen formed at the Pd catalyst is possible even in the presence of H2O2. No

309

significant changes in total dissolved iron concentration were observed after the ACF reaction.

310

Thus, we suggest that iron switches between its oxidation states via reactions 1 and 9 with

311

[Fe H2 O)5 OH)]

312

hydrolysis equilibria of FeII and FeIII, respectively.17 Due to their low formation constants17, FeIII-

313

peroxide complexes are of minor importance in terms of species concentrations under the

314

applied conditions (see SI, part S2).

315

H2, which was generally provided as 80 mL gas-phase reservoir over 50 mL reaction suspension

316

(i.e. 67 mmol H2/L water), was 45% consumed within 60 min reaction under the conditions of

317

the ACF experiment shown in Fig. 2. In the ‘breathing system’, N2 was sucked into the vessel.

318

O2 was detected only in trace amounts within the gas phase (≤ 4 vol-%) during the whole

319

reaction period. This can be due to two reasons: I) the H2O2 conversion to oxygen in the

320

regular Fenton cycle (with Eq. 2 as initial step) is largely eliminated in the ACF system, and II)

321

intermediately formed oxygen is rapidly hydrogenated.

322

OH consumers. However, in a heterogeneous system, several of these parameters are

2+

and [Fe H2 O)6 ]

2+

as the dominant species based on the pH-dependent

3.3. Indications of reactive species responsible for MTBE oxidation

323

The typical intermediates known from OH-radical-driven MTBE oxidation, i.e. TBA, TBF and

324

acetone32,33,35,39, also appeared in the ACF system, indicating that MTBE degradation follows a

325

similar pathway to that in other AOP systems (see SI, Fig. S4.1 and discussion).



Page 12 of 23

326

TBA, which is known to act as a scavenger of ●OH, largely inhibited MTBE degradation when

327

added to the ACF system at high surplus (SI, Fig. S4.2). In addition, compound-specific stable

328

isotope analysis (CSIA) was applied. Isotope fractionation effects are specific for a certain

329

mechanism and transition state of the reaction.40 In previous studies, CSIA with H/D and

330

13

331

bly FeIV) in photo-Fenton41 and Fe-zeolite-catalyzed Fenton reactions42 or indicated similarity of

332

reactive species (i.e. ●OH) in orthoferrite-catalyzed and homogeneous Fenton systems43. Based

333

on competition kinetics, the slope of the ln-ln plot of the residual concentrations of two sub-

334

strates in the same reaction system, such as MTBE-d0 and -d12, yields AKIEH/D, the so-called

335

apparent (or observable) kinetic deuterium isotope effect (Eq. 18). This value is the ratio of the

336

second-order rate constants for the first attack in the oxidation of the two isotopologues

C/12C compound isotopologues revealed the occurrence of additional oxidants to ●OH (possi-

337

(kH,OH/kD,OH), which is assumed to be H-abstraction from one of the methyl groups of MTBE by

338

●

OH.

ln

-6 -9 =6,>6 / ln = :;6

Eq. 18

339

From the competition kinetics plot, an associated AKIEH/D = 2.01 ± 0.04 was obtained for

340

MTBE-d0 and -d12 degradation in the ACF system (Fig. 3). The excellent conformity with the

341

AKIEH/D observed in a simple stoichiometric Fenton system (2.00 ± 0.04) is a strong argument

342

for ●OH also being the dominant reactive species in the ACF system.

Figure 3: Competition kinetics plot for degradation of MTBE-d0 and -d12 by ACF (●, CFe = 18 µM, CPd = 5 mg L-1, pH = 3, CH2O2,0 = 15 mM, Vwater = 50 mL, Vgas = 80 mL (100 vol-% H2), CMTBE-d0,0 = CMTBE-d12,0 = 0.2 mM) and a stoichiometric Fenton reaction (, CFe = 15 mM, CH2O2,0 = 15 mM, CMTBE-d0,0 = CMTBE-d12,0 = 0.2 mM, pH = 3, after 4 min of reaction).


Page 13 of 23


343

3.4. Effects of pH, dissolved iron, H2O2 and H2 concentration on the reaction kinetics

344

For MTBE degradation in the ACF system we observed a pH optimum similar to that generally

345

known for CF reactions (Fig. S5a). MTBE turnover is highest at pH = 3, significantly lower at pH

346

2 and pH 4, and largely inhibited if the solution pH approaches pH 7. This is most likely due to

347

the unfavorable speciation of FeIII which forms insoluble oxyhydroxide precipitates at pH ≥ 3. In

348

addition, H2O2 decomposes faster if the pH is increased from pH = 3 to 7 (Fig. S5a). This effect

349

can be ascribed mainly to the parasitic decomposition and hydrogenation of H2O2 at the Pd

350

catalyst which (in the absence of iron) becomes faster with pH increasing from 3 to 6 (Fig. S5b),

351

in conformity with results of Choudhary et al.19 Thus, even though precipitation of FeIII at pH ≥

352

3 could be counteracted in the ACF system by addition of suitable complexing agents, these

353

efforts would be undermined by the pH effect on parasitic H2O2 decomposition by Pd, which is

354

detrimental for conducting Pd-based ACF processes at near-neutral pH. Thus, the optimal pH

355

for the ACF is pH = 3.

356

The importance of the various ●OH consumers in a reaction system can be estimated from the

357

products of their concentration and rate constants for reaction with ●OH. Even though con-

358

sumption rates for ●OH at the Pd surface cannot be estimated, the importance of the bulk-

359

phase consumers FeII, MTBE and H2O2 can be compared for various conditions. Under the con-

360

ditions applied (CFeII,0 ≤ 36 µM, CMTBE,0 = 0.17 mM), iron is not relevant as consumer of ●OH since

361

it is outcompeted by MTBE with CFeII,0 × k●OH,FeII = 0.04 × (CMTBE,0 × k●OH,MTBE). Thus, the effect of

362

an increased iron concentration on ●OH formation via the Fe redox cycle (Eqs. 1 and 9) is

363

predominantly positive, as shown by the correlation between k’MTBE,initial and CFe in a range of

364

0.5 to 2 mg L-1 (Fig. 4a). Even though higher amounts of dissolved iron increase the rate of

365

contaminant degradation, 18 µM (1 mg L-1) is selected as optimal for the ACF in order to

366

comply with discharge limits and avoid the need of post-treatment for iron removal.

367

Increasing the H2O2 concentration in Fenton-type reactions is only useful up to the point at

368

which H2O2 becomes the dominant quencher of

369

concentration will not increase contaminant degradation rates because the steady-state

370

concentration of ●OH is not increased. MTBE and H2O2 are both equally relevant as consumers

371

of ●OH for CMTBE,0 = 0.17 mM and CH2O2 = 7.5 mM. Thus, we estimate an optimal H2O2

372

concentration below 30 mM, where the ●OH quenching ratio of MTBE : H2O2 is 1 : 3. As

373

discussed in section 3.2, H2O2 can also play a detrimental role as consumer of activated

374

hydrogen at the Pd surface, thus competing with FeIII reduction. The influence of H2O2

375

concentration on the ACF system is illustrated in Fig. 4b, based on MTBE turnover within 60

376

min of reaction. There is clearly an optimal range of H2O2 concentration. In line with the

377

increasing role as consumer of ●OH, H2O2 concentrations ≥ 30 mM are detrimental for MTBE

378

degradation. MTBE turnover was almost equal, and thus in the optimal range, when reactions

●

OH. Any further increase in H2O2



Page 14 of 23

379

were started with initial H2O2 concentrations of 1.7 to 15 mM. Obviously, reactions with low

380

initial H2O2 concentration demand more frequent re-dosing of H2O2 (SI, Fig. S6).

381

The rather low impact of CH2O2 on MTBE degradation is in line with the finding that pseudo-first

382

order kinetics applies for MTBE degradation if started with CH2O2,0 = 15 mM (Fig. 2a), even

383

though H2O2 concentration declines to C/C0 = 0.1 in 60 min. The relative insensitivity of

384

contaminant degradation rate towards CH2O2 over a wide concentration range supports the

385

hypothesis that the rate-determining step of the Fenton cycle, i.e. FeIII reduction, is performed

386

at the Pd catalyst and is thus not dependent on H2O2. In conclusion, H2O2 concentration is in

387

the optimal range if it is: I) high enough to utilize the FeII production rate (via Eq. 9) and thus

388

maximize the rate of ●OH production (Eq. 1), and II) lower than the ●OH quenching limit as

389

described above.

390

The minimum molar ratio of H2O2 : MTBE for achieving ≥ 95 % degradation within 60 min is 59,

391

which is at the higher end of the range of 10 to 55 reported for conventional homogeneous

392

Fenton systems for MTBE degradation in other studies.31-36 However, as discussed before,

393

conventional Fenton systems require significantly higher Fe doses (10–260 mg L-1) and need

394

post-treatment for iron removal.

395

For evaluating the efficiency of H2O2 utilization the initial reaction period (until 25 to 50% of

396

target contaminant conversion) is more appropriate, since at higher turnover the relative

397

contribution of other ●OH consumers such as degradation intermediates becomes increasingly

398

relevant. Thus, we determined the utilization efficiency η50 which is mol of MTBE degraded per

399

mol of H2O2 consumed (in %, Eqs. 16 and 17) in the reaction period until 50% MTBE conversion

400

are reached. Among the experiments with various modes of H2O2 dosage (Figure 4b) which

401

reached comparable MTBE conversion rates (≥ 95% degradation in 60 min), the highest η50

402

value of 3% was obtained with the lowest stationary H2O2 concentration (1.7 mM H2O2 dosed

403

every 10 min), while η50 = 1.4% for experiments started with 15 mM H2O2. Taking into account

404

that non-productive consumption of H2O2 at the Pd catalyst was estimated to reduce E1 to

405

≤70%, the contribution of E2, i.e. consumption of ●OH by other species than MTBE, appears to

406

be the critical factor for H2O2 utilization efficiency in the ACF system. Relevant ●OH consumers

407

beside MTBE are its oxidation intermediates, H2O2 (which is largely ruled out at the lowest

408

applied H2O2 concentration in Fig. 4), but also activated hydrogen at the Pd surface. Dissolved

409

molecular hydrogen (H2) can be neglected as ●OH consumer because of its low aqueous

410

solubility (0.8 mM) and its low reactivity (kH2,OH ≈ 5 x 107 L mol-1 s-1).

411

For comparison, η50 = 9.2% can be calculated from the data reported by Hwang et al.34 for

412

MTBE degradation in a homogeneous Fenton reaction (k’MTBE = 9.0 h-1, k’H2O2 = 1.4 h-1, C0,MTBE =

413

0.11 mM, C0,H2O2 = 6 mM, CFe(NO3)3 = 5 mM, pH 3). Heterogeneous Fenton-like catalysts, such as

414

iron oxides2,44, iron containing perovskites43 or iron supported on alumosilicates44, even though 14 ACS Paragon Plus Environment

Page 15 of 23


415

they are able to work at near neutral pH, show much lower catalytic activities (contaminant

416

half-lifes in the range of hours at g L-1 solid catalyst and H2O2 concentrations) and lower H2O2

417

utilization efficiencies (for iron oxides typically > L●OH, appears to be a prerequisite for the

504

efficiency of the ACF system, since it shifts the location of ●OH production and consumption

505

into the water phase, i.e. towards dissolved contaminants and away from the solid catalyst

506

surface, where active hydrogen species can quench ●OH by recombination. Nevertheless, this

507

●

508

H2O2 for contaminant degradation in the ACF system. Internal mass-transfer limitations within

509

the microporous particles of the commercial Pd on γ-Al2O3 catalyst cannot be excluded for the

510

fast reactions involved in the ACF system, which can also affect oxidant efficiency. Thus, future

511

studies should include optimization of the Pd catalyst support including the use of

512

nanoparticles instead of microporous supports, such that FeII can unhindered diffuse into the

513

bulk water phase, away from ●OH quenching surfaces. In summary, the ACF system allows the

514

excellent activity of dissolved FeII to be utilized for ●OH production at extremely low catalyst

515

concentration, thus eliminating the need to remove or recycle the iron catalyst. This is

516

achieved by introducing a clean reductant (H2) and a highly efficient, recyclable solid Pd

517

catalyst for accelerating the slowest step of the Fenton cycle, i.e. FeIII reduction. In contrast to

518

previous approaches for enhancement of Fenton reactions by addition of reductants such as

519

hydroxyl amine14 or ascorbic acid15, the reductant H2 applied in ACF is inexpensive, leaves no

520

residues in the treated water and is less reactive towards oxidation by ●OH (by up to two

521

orders of magnitude) than many water pollutants. Using externally supplied H2O2 and low

522

concentrations of H2 (below the level of explosive gas), the ACF system facilitates efficient and

523

safe contaminant oxidation with only trace amounts of iron. Thus, the ACF approach presented

524

in this study offers potential for the development of significantly improved water treatment

525

technologies.

526

Supporting Information: Details on experimental setup, impact of Pd (without H2) in catalytic

527

Fenton reactions, experimental details and results on iron speciation in ACF, MTBE oxidation

528

intermediates, quenching experiments, influence of pH on ACF and H2O2 decomposition by

529

Pd/H2, ACF with various modes of H2O2 dosage, in-situ formation of H2O2, ACF at reduced H2

530

concentration, stability of Pd/Al2O3 catalyst.

OH quenching process certainly remains an important factor for the utilization efficiency of



Page 20 of 23

531

Acknowledgement:

532

Funding by ESF (Grant: 24127008/TG74) and Sächsische Aufbaubank is gratefully

533

acknowledged.

534

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