Degradation of Bisphenol A by Peroxymonosulfate Catalytically

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Degradation of Bisphenol A by Peroxymonosulfate Catalytically Activated with Mn1.8Fe1.2O4 Nanospheres: Synergism between Mn and Fe Gui-Xiang Huang, Chu-Ya Wang, Chuan-Wang Yang, Pu-Can Guo, and Han-Qing Yu Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.7b03007 • Publication Date (Web): 06 Oct 2017 Downloaded from http://pubs.acs.org on October 6, 2017

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Environmental Science & Technology

Degradation of Bisphenol A by Peroxymonosulfate Catalytically Activated with Mn1.8Fe1.2O4 Nanospheres: Synergism between Mn and Fe

Gui-Xiang Huang, Chu-Ya Wang, Chuan-Wang Yang, Pu-Can Guo, Han-Qing Yu* CAS Key Laboratory of Urban Pollutant Conversion, Department of Chemistry, University of Science & Technology of China, Hefei, 230026, China

1

ACS Paragon Plus Environment


1

ABSTRACT

2

A high-efficient, low-cost, and eco-friendly catalyst is highly desired to activate

3

peroxides for environmental remediation. Due to the potential synergistic effect

4

between bimetallic oxides’ two different metal cations, these oxides exhibit superior

5

performance in the catalytic activation of peroxymonosulfate (PMS). In this work,

6

novel Mn1.8Fe1.2O4 nanospheres were synthesized and used to activate PMS for the

7

degradation of bisphenol A (BPA), a typical refractory pollutant. The catalytic

8

performance of the Mn1.8Fe1.2O4 nanospheres was substantially greater than that of the

9

Mn/Fe monometallic oxides and remained efficient in a wide pH range from 4 to 10.

10

More importantly, a synergistic effect between solid-state Mn and Fe was identified in

11

control experiments with Mn3O4 and Fe3O4. Mn was inferred to be the primary active

12

site in the surface of the Mn1.8Fe1.2O4 nanospheres, while Fe(III) was found to play a

13

key role in the synergism with Mn by acting as the main adsorption site for the

14

reaction substrates. Both sulfate and hydroxyl radicals were generated in the PMS

15

activation process. The intermediates of BPA degradation were identified and the

16

degradation pathways were proposed. This work is expected to help to elucidate the

17

rational design and efficient synthesis of bimetallic materials for PMS activation.

2


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18

INTRODUCTION

19 20

Sulfate radical (SO4•−)-based advanced oxidation processes have received increasing

21

attention for their applications in the field of environmental protection, including the

22

degradation of recalcitrant organics in water,1 disinfection,2 and the disintegration of

23

activated sludge.3 Compared with hydroxyl radical (•OH), SO4•− possesses several

24

advantages including a more positive reduction potential of 2.5-3.1 V (vs. 1.8-2.7 V

25

for •OH), 4 a pH-independent reactivity,1 a higher oxidation selectivity5 and a longer

26

lifetime (t1/2 = 30-40 µs, vs. 10−3 µs for •OH).6 In a typical process, SO4•− is generated

27

by activating peroxymonosulfate (PMS) catalytically with various transition-metal

28

catalysts.1 These catalysts include such metals as Co, Ag, Cu, Mn and Fe; although Co,

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Ag and Cu have been identified as excellent PMS activators, they are practically

30

limited by their relatively high toxicity and low geological reserves.7 Therefore, the

31

development of efficient Fe- and Mn-based catalysts becomes a priority for PMS

32

activation.

33

Commonly used Fe-based materials include zero valent iron (ZVI), Fe3O4 and

34

Fe2O3. However, the conversion of Fe0 to Fe3+ in the catalytic reaction can deactivate

35

ZVI when the catalyst is reused, and the catalytic performances of pure Fe3O4 and

36

Fe2O3 are generally low.1 Wang and co-workers have investigated the performance of

37

various monometallic Mn oxides and the factors that govern their catalytic

38

activities.8-14 It has been demonstrated that most Mn oxides have catalytically

39

activated PMS well, but at the cost of a high dose of PMS (Table S1). Thus, an

40

effective strategy should be pursued to further enhance the performance of Fe/Mn

41

oxides.

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The synthesis of bimetallic oxides is recognized an effective method to improve

43

the catalytic activity of materials in both energy- and environment-related

44

applications.15-18 It was reported that CuFeO2 exhibited a higher reactivity and

45

stability than Cu2O and Fe2O3, and a synergistic catalytic effect between solid Cu(I)

46

and Fe(III) was identified to be attributed to the accelerated reduction of Fe(III).18 A

47

synergistic effect in some Mn-Fe bimetallic oxides has been reported.19-21 However, it 3



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48

was simply deduced from the difference in the catalytic performances of the

49

bimetallic oxide and the two corresponding monometallic oxides, without even

50

normalizing for their specific surface areas. Moreover, the synergistic mechanism,

51

especially the role of Fe in Mn-Fe bimetallic systems, remains unknown and deserves

52

further investigations.

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Therefore, in this work, novel Mn-Fe bimetallic oxide Mn1.8Fe1.2O4 (hereinafter

54

abbreviated as MnFeO) nanospheres were synthesized by heating the designed

55

precursor, a Mn-Fe Prussian blue analogue (PBA), in air. The as-prepared MnFeO

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nanospheres were used to activate PMS to degrade bisphenol A (BPA), a widespread

57

endocrine-disrupting pollutant in the aquatic environment.22,23 The catalytic

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performance of the MnFeO nanospheres was examined in details and the catalytic

59

mechanism was also investigated. More importantly, the synergistic effect was

60

explored in order to understand the influence between Mn and Fe in the MnFeO

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nanospheres. Additionally, the degradation pathways of BPA were established based

62

on the identified intermediates and the stability of the MnFeO nanospheres was also

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evaluated. This work is expected to provide useful information for the further

64

development of advanced catalysts for sulfate radical-based advanced oxidation

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processes.

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EXPERIMENTAL SECTION

68 69

Chemicals and Reagents. Unless otherwise specified, all chemicals and reagents

70

were

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(2KHSO5·KHSO4·K2SO4, 4.5% active oxygen) was purchased from Beijing J&K Co.,

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China. BPA and α-Fe2O3 nanoparticles (30 nm) were purchased from Aladdin Co.,

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China. Acetonitrile, methanol (gradient grade) and 5,5-dimethyl-1-pyrroline-N-oxide

74

(DMPO) were purchased from Sigma-Aldrich Co., China. Other reagents were

75

purchased from Shanghai Chemical Reagent Co., China.

of

analytical

grade

and

used

without

further

purification.

PMS

76

Synthesis of MnFeO Nanospheres and Mn3O4. The MnFeO nanospheres were

77

synthesized according to a modified protocol reported previously.24 In brief, 4


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MnCl2·4H2O (6.25 mmol) and polyvinylpyrrolidone (PVP, 0.75 g) were first

79

dissolved in 25 mL deionized water. An aqueous solution of K3[Fe(CN)6] (125 mM,

80

25 mL) was then poured quickly into the aforementioned solution with vigorous

81

stirring. The obtained colloid solution was stirred for 30 min further and then aged for

82

1 day. The resulting khaki precipitate was collected via centrifugation and washed

83

with distilled water and ethanol several times. The product was then dried at 70 °C for

84

24 h in a vacuum oven. To prepare the MnFeO nanospheres, the obtained solid

85

powder was heated to 400 °C with a temperate ramp of 2 °C min–1 and kept at the

86

same temperature for 1 h in air. Mn3O4 was synthesized using a modiﬁed

87

hydrothermal method reported previously.25

88

Characterization of the Catalysts. The morphological and textural properties of

89

the catalysts were examined with a field emission scanning electron microscope

90

(SEM) (JSM-6700F, JEOL Co., Japan) and a transmission electron microscope (TEM)

91

(H7650, Hitachi Co., Japan). The specific surface areas of the catalysts were

92

measured using the Brunauer-Emmett-Teller (BET) N2 adsorption-desorption method

93

with a Builder 4200 instrument (Tristar II 3020M, Micromeritics Co., USA). The

94

X-ray powder diffraction (XRD) patterns of the samples were obtained using a Philips

95

X’Pert PRO SUPER diffractometer equipped with graphite monochromatized Cu Kα

96

radiation (λ = 1.541874 Å). The MnFeO nanospheres were dissolved in hydrochloric

97

acid (12 M) so that their element composition could be analyzed using an inductively

98

coupled plasma-mass spectrometer (ICP-MS) (PlasmaQuad 3, Thermo Fisher Inc.,

99

USA). The valence states of the constituent elements were determined using X-ray

100

photoelectron spectroscopy (XPS) (ESCALAB250, Thermo Fisher Inc., USA), and

101

the binding energy was calibrated with the C 1s peak at 284.8 eV. The surface

102

properties of the MnFeO samples before and after the degradation reaction were

103

characterized using Raman spectroscopy (inVia confocal Raman, Renishaw Co., UK)

104

and Fourier transform infrared spectroscopy (FTIR, Vertex 70, Bruker Co., Germany).

105

BPA Degradation Experiments. Unless otherwise specified, all of the

106

degradation experiments were carried out in a 100-mL reactor containing 40 mL of

107

BPA solution (10 mg L–1) at room temperature (25 ± 2 °C); the pH values of the 5



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108

reaction solutions were adjusted with 0.1 M NaOH or H2SO4 and buffered with borate

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when needed, which was recorded with an pH meter (model PHS-3E, INESA Co.,

110

China). Typically, the catalyst of 4 mg was added to 40-mL BPA solution. After 1-min

111

ultrasonic dispersion, a uniform suspension was created, which was then stirred for 15

112

min to establish the adsorption–desorption equilibrium. Then, PMS of 8 mg was

113

added to the suspension to initiate the reaction. One milliliter of the suspension was

114

withdrawn and quenched with half a milliliter of ethanol at given time intervals. The

115

sample was centrifuged immediately to separate the solid and liquid, and the

116

supernatant was collected for subsequent BPA concentration measurements, which

117

were made within 2 h. For recyclability tests, the catalyst was recovered by

118

centrifuging the sample, washing the catalyst several times with distilled water, and

119

drying it at 40 °C in a vacuum oven. All experiments were carried out in duplicate or

120

triplicate, and the average data with their standard deviations are presented.

121

Analysis. The BPA concentration was analyzed using high-performance liquid

122

chromatography (HPLC) (LC-16, Shimadzu Co., Japan) with a C18 column. An

123

acetonitrile /water (containing 0.1% formic acid) mixture (50:50, v/v) was used as the

124

mobile phase at a ﬂow rate of 0.5 mL min–1, and the detection wavelength was 273

125

nm. The total organic carbon (TOC) concentration was measured using a TOC

126

analyzer (Muti N/C 2100, Analytik Jena AG, Germany). Free radicals were detected

127

using an electron paramagnetic resonance (EPR) spectrometer (JES-FA200, JEOL Co.,

128

Japan). The intermediate products of BPA degradation were determined using a gas

129

chromatography-mass spectrometer (GC-MS) (Agilent Co., USA) with an Agilent

130

7890B GC system in combination with an Agilent 5977B single quadrupole mass

131

spectrometric detector. The concentrations of leached manganese and iron were

132

measured using an ICP-MS.

133 134

The BPA degradation kinetics were fit by the pseudo first-order model and the apparent rate constant (k) was calculated according to eq 1:26 ln( ∕ ) = −

(1)

135

where Ct is the BPA concentration at a certain reaction time (t) and C0 is the initial

136

BPA concentration. 6


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137 138

RESULTS AND DISCUSSION

139 140

Physicochemical Characteristics of MnFeO Nanospheres. The SEM and TEM

141

images clearly show the sphere-like morphology of the as-synthesized product with a

142

diameter ranging from 100 to 500 nm (Figure 1a). The BET surface area of the

143

product was 58 m2 g−1, determined by a N2 adsorption–desorption measurement

144

(Figure S1). The XRD analysis was applied to confirm the crystallographic structure

145

and phase purity of the product, and the result shows that all the characteristic

146

diffraction peaks were identical to those of spinel Fe3O4 (JCPDS card No. 75-0449)

147

(Figure 1b). No other crystalline impurities were detected, indicating the single phase

148

of the product. The element composition of the product was analyzed using ICP-MS,

149

and the calculated atomic ratio of transition metals was approximately 1.5:1 (Mn:Fe).

150

As a result, the chemical formula of the product was designated Mn1.8Fe1.2O4. In

151

addition, the phase of the as-prepared Mn3O4 and commercial Fe3O4 and Fe2O3 were

152

also confirmed by analyzing their XRD patterns (Figure S2).

153

Catalytic Performance of MnFeO Nanospheres in PMS Activation for BPA

154

Degradation. The catalytic activity of the MnFeO nanospheres was evaluated by

155

activating PMS to degrade BPA. Since the solution pH would drop to a certain degree

156

in the reaction without any buffers (Figure S3), borate buffer of 20 mM was dosed to

157

control the solution pH when needed (Figure S4), which had limited influence on the

158

catalytic performance of the MnFeO nanospheres (Figure S5). As shown in Figure 2a,

159

when PMS or MnFeO was used alone, only less than 6% of BPA was degraded,

160

indicating that both the intrinsic oxidizing power of PMS and the adsorption capacity

161

of the MnFeO nanospheres were negligible under the tested conditions. However,

162

when the MnFeO nanospheres and PMS were used together, more than 95% of BPA

163

was degraded within 30 min, which was much higher than those of the commercial

164

Fe3O4 (3%) and the synthetic Mn3O4 (23%). Thus, the catalytic ability of the MnFeO

165

nanospheres was much greater than those of Fe3O4 and Mn3O4. This was further

166

confirmed by comparing the BPA removal efficiencies after their normalization by the 7



167

specific surface areas (Figure S6). In the homogeneous control tests (Figure S7),

168

negligible BPA was degraded in the Mn2+ of 20 µg L–1 (as shown by the ICP-MS

169

result after the reaction of the PMS/MnFeO system, Table S2) and the leaching

170

solution control groups, indicating that the BPA degradation mainly occurred at the

171

surface of the MnFeO nanospheres through a series of heterogeneous catalytic

172

reactions. As shown in Figure 2b, the BPA degradation kinetics were well fit by the

173

pseudo first-order model and the apparent rate constant of the MnFeO nanospheres

174

was estimated to be 0.10 min–1.

175

The impact of the solution pH on the BPA degradation was examined and

176

negligible changes in the removal efficiency were observed in a wide pH range from

177

4.2 to 10.2 (Figure 2c). Previous studies have shown that the solution pH influenced

178

the catalytic behaviors of heterogeneous catalysts from different aspects, e.g., the

179

ionization of PMS and pollutant molecules, the surface charges of the catalysts, the

180

transformation from SO4•− to •OH and their oxidation potentials.27 Under the acidic

181

condition, more positive charges at the MnFeO surface would reinforce the affinity

182

for PMS, which existed mainly as HSO5− according to its pKa values (pKa1 < 0, pKa2

183

= 9.4),28 but the adverse effect also existed due to the stabilization effect of H+ on

184

HSO5−.29 At pH 10.2, the catalyst surface was negatively charged as the pHpzc (pH of

185

point of zero charge) of the substituted magnetites was around 6.8,30 which was

186

unfavorable for the absorption of HSO5−, SO52− and BPA anions. However, a positive

187

factor is that the increasing amount of surface hydroxyls could also accelerate the

188

decomposition of PMS.31 Therefore, the impact of the solution pH on BPA

189

degradation came from the integrative actions of various changes in the

190

physicochemical properties of all the substances involved. The high efficiency of the

191

PMS/MnFeO oxidation process under acidic, neutral and alkaline conditions suggests

192

its promising potential for the treatment of various wastewaters.

193

As shown in Figure 2d, the BPA degradation efficiency exhibited a positive

194

dependence on the MnFeO nanospheres dosage. When the dosage was increased to

195

0.2 g L–1, complete degradation of BPA was achieved within 15 min only. In the case

196

of removing organic pollutants with a comparable PMS dosage, the catalytic 8


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197

performance of the MnFeO nanospheres was much more robust than that of most of

198

the Mn/Fe-oxide catalysts reported previously (Table S1). To ensure both the nearly

199

complete BPA degradation and the modest reaction kinetics, a medium dosage, i.e.,

200

0.1 g L–1, was used in the subsequent experiments. As shown in Figure S8,

201

approximately 80% of the TOC was removed within 30 min at the catalyst dosage of

202

0.1 g L–1, indicating the excellent mineralization efficiency of the PMS/MnFeO

203

process.

204

Additionally, the stability of the MnFeO nanospheres was also evaluated (Figure

205

S9). Although the BPA degradation efficiency decreased obviously after the first run

206

(when the catalyst was washed simply by water), it was able to fully recover after the

207

thermal treatment at 400 °C for 1 h in air. The Raman and FTIR spectra of the MnFeO

208

samples before and after the reaction were used to characterize the changes in the

209

catalyst surface. The Raman spectra show that no phase change occurred during the

210

reaction (Figure S10). In the FTIR spectra of the sample after the reaction (Figure

211

S11), some bands emerge at 1500 and 1460 cm−1 and at 1217 and 1175 cm−1, which

212

could be assigned to the skeletal vibrations of the aromatic rings and the bending

213

vibrations of aromatic C-H, respectively.32-35 This result suggests that the aromatic

214

intermediates deposited on the catalyst surface led to the deactivation of the MnFeO

215

nanospheres, and they could be removed effectively by thermal treatment. Moreover,

216

the metal leaching properties of the catalyst were also investigated. As shown in

217

Table S2, the concentration of leaching Mn and Fe ions after the reaction at various

218

pHs in Figure 2c and after different cycles in Figure S9 were all below 1 mg L–1,

219

further demonstrating the stability of the MnFeO nanospheres.

220

Synergistic Effect between Mn and Fe on Catalytic Performance. To explore

221

the potential synergistic effect between Mn and Fe in the MnFeO nanospheres, a

222

series of combinations of Fe3O4 and Mn3O4 were used in the control group. With the

223

total dosage of catalysts being kept constant (0.1 g L–1), the catalytic performance of

224

the mixture of Fe3O4 and Mn3O4 was supposed to be better than that of pure Fe3O4 but

225

worse than that of pure Mn3O4. However, Figure 3 reveals that the three

226

combinations of 1:1 to 1:3 (Fe3O4:Mn3O4, w:w) performed better than both pure 9



227

Fe3O4 and pure Mn3O4, demonstrating that there was a synergistic effect between Mn

228

and Fe oxides in the activation of PMS. It is worth noting that the best performance

229

was achieved when the weight ratio of Fe3O4 and Mn3O4 was 1:1.5 and the molar

230

ratio of Fe to Mn was also 1:1.5, which is the same as that of the as-prepared MnFeO

231

nanospheres. However, even so, the catalytic activity of the mixture was still lower

232

than that of the MnFeO nanospheres (Figure 2a). As shown in Figure S12, the

233

effective contact between Fe and Mn mainly came from the wrapping of the Fe3O4

234

particles by the Mn3O4 nanowires, which was much less compact than that in the

235

lattice of the MnFeO nanospheres. Therefore, the catalytic performance of the

236

mixtures of Fe3O4 and Mn3O4 was substantially poorer than that of MnFeO, indicating

237

that the sufficient and efficient contact between Mn and Fe is essential for their

238

synergism.

239

Radicals in the PMS/MnFeO Nanospheres System. To identify the radical

240

species involved in the BPA degradation, EPR experiments using DMPO as the

241

spin-trapping agent were carried out. It is commonly accepted that both SO4•− and

242

•OH can be formed during the catalytic activation of PMS by transition metal

243

oxides.18, 28, 36 As shown in Figure 4a, no peaks were identified in the test groups of

244

the PMS and PMS+BPA solutions, indicating that no radicals were produced in the

245

absence of the MnFeO nanospheres. When the MnFeO nanospheres were added, a set

246

of peaks were obtained and could be assigned to DMPO•-OH (with hyperfine

247

couplings αN = αβ-H = 14.9 G) and DMPO•-SO4− (with hyperfine splitting constants of

248

αN = 13.2 G, αβ-H = 9.6 G, αγ-H1 = 1.48 G and αγ-H2 = 0.78 G).12, 37, 38 In the presence of

249

BPA, the signal intensities of both DMPO•-OH and DMPO•-SO4− decreased

250

substantially compared with the control group without BPA, indicating that both •OH

251

and SO4•− were able to react with BPA, resulting in its degradation.

252

To further confirm the contributions of the two radicals, ethanol (EtOH) and

253

tert-butyl alcohol (TBA) were used as radical scavengers. EtOH possesses a high

254

reactivity with both •OH (• = 1.9 × 109 M s ) and SO4•− (• = 1.6 ×

255

107 M s ), and TBA has a good reactivity with •OH (• = 6.0 × 108 M s) 10


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256

but poor reactivity with SO4•− (• = 4.0 × 105 M s ).39, 40 Figure 4b shows

257

that EtOH obviously inhibited the BPA degradation and such an inhibition was

258

enhanced with the increasing scavenger dosage. This could be attributed to the

259

competitive consumption of •OH and SO4•− by EtOH. However, whether at a low (1.2

260

M) or high (6.0 M) concentration, TBA always inhibited BPA degradation more than

261

EtOH did, which is inconsistent with the most results reported.18, 38, 41, 42 This result is

262

probably because of the masking effect on the bonding sites dispersed in the surface

263

of the MnFeO nanospheres caused by the high viscosity of TBA.31, 43 It should also be

264

noted that the negligible impact of alcohols at 1.2 M, which is four orders of

265

magnitude more than that of BPA, on the BPA degradation kinetics is an unusual

266

phenomenon. To find out the reason for this observation, we conducted additional

267

EPR tests by adding different levels of EtOH into the reaction solution. As shown in

268

Figure S13, when 1.2 M EtOH was added, the peaks of both SO4•− and •OH only

269

weakened slightly in the intensity, and they were still recognizable in the presence of

270

12 M EtOH. Similar result has been reported previously.44 It should also be noted that,

271

in the presence of BPA, the signal intensities of the radicals decreased more

272

substantially than in the case of 1.2 M EtOH. These results indicate that, on the one

273

hand, the generated radicals mainly adhered to the catalyst surface, which could

274

hardly be cleaned up by alcohols, even at an ultrahigh concentration; on the other

275

hand, the MnFeO surface had a stronger affinity for BPA than alcohols, due to the

276

stronger coordination of metals with phenolic hydroxyls than alcoholic hydroxyls.

277

Mechanism for the Generation of Radicals and the Synergistic Effect. The

278

XPS spectra of the MnFeO nanospheres before and after the reaction were used to

279

explore the PMS activation mechanism. The binding energies and relative intensities

280

are summarized in Table 1 based on the deconvolution of Mn 2p and Fe 2p XPS

281

spectra (Figure S14). The Mn 2p spectrum (Figure S14a) was composed of a

282

spin-orbit doublet of Mn 2p1/2 and Mn 2p3/2 with a binding energy gap of 11.5 ± 0.1

283

eV, and the deconvoluted Mn 2p3/2 spectrum displayed four peaks with binding

284

energies at 640.8, 641.9, 643.0 and 644.5 eV, which could be assigned to Mn(II), 11



Page 12 of 31

285

Mn(III), Mn(IV), and the shake-up peak, respectively.45 After the catalytic reaction,

286

the relative intensity of Mn(IV) remained constant, and only 1% of the total Mn was

287

transformed from Mn(II) to Mn(III). For the Fe 2p spectrum, it was reported that

288

Fe(III) and Fe(II) in octahedral sites of magnetite were distinguishable in the XPS

289

analysis because the core-hole lifetime in the photoemission process is much shorter

290

than the electron hopping frequency by approximately four orders of magnitude.46

291

Hence, the deconvoluted Fe 2p3/2 spectrum (Figure S14b) displayed three peaks with

292

binding energies at 710.2, 711.0 and 713.0 eV, which could be assigned to octahedral

293

Fe(II), octahedral Fe(III), and tetrahedral Fe(III), respectively.47, 48 After the catalytic

294

reaction, the relative intensity of tetrahedral Fe(III) remained constant, but 7% of the

295

total Fe was transformed from octahedral Fe(II) to octahedral Fe(III).

296

Therefore, these results suggest that the activation of PMS occurred on the

297

catalyst surface. Both Mn(II) and Fe(II) donated electrons to PMS and thus initiated

298

its decomposition (eqs. 2 and 3). Meanwhile, Mn(III) also activated PMS through an

299

additional one-electron donation (eq. 4).10 Mn and Fe at higher valence states were

300

then reduced by HSO5− to complete the redox cycle (eqs. 5 to 7),1 which made the

301

catalytic action of the MnFeO nanospheres work continuously. In this process, •OH

302

was generated through the reaction between SO4•− and H2O/OH− (eqs. 8 and 9).40 The

303

generated SO4•− and •OH attacked BPA through a series of reactions including

304

electron transfer, electrophilic/radical addition and hydrogen abstraction,5, 49 which

305

decomposed BPA into various intermediates and finally mineralized into CO2 and

306

H2O (eq. 10). • ≡Fe(II) + HSO ! → ≡Fe(III) + SO# + OH

(2)

• ≡Mn(II) + HSO ! → ≡Mn(III) + SO# + OH

(3)

• ≡Mn(III) + HSO ! → ≡Mn(IV) + SO# + OH

(4)

• ( ≡Fe(III) + HSO ! → ≡Fe(II) + SO! + H

(5)

• ( ≡Mn(IV) + HSO ! → ≡Mn(III) + SO! + H

(6)

• ( ≡Mn(III) + HSO ! → ≡Mn(II) + SO! + H

(7)

( SO •# + H, O → SO 2 # + •OH + H , < 6 × 10 M s

(8)

12


Page 13 of 31


2 SO •# + OH → SO 2 # + •OH, = 6.5 × 10 M s

(9)

SO •# /•OH + BPA → intermediates → CO, + H, O 307

(10)

In previous works, an electron transfer between Mn and Fe was proposed to be

308

responsible for the synergistic mechanism.50,

309

potentials of the metals (eqs. 11 and 12), the reduction of Mn(III) by Fe(II) is

310

thermodynamically favorable (eq. 13).13 However, with a consideration of the

311

reduction potentials of HSO5−/SO4•− (2.5-3.1 V) and HSO5−/SO5•− (1.1 V),42, 52 the

312

regeneration of Fe(II) is the rate-determining step in the activation of PMS. For the

313

Mn(III)/Mn(II) redox pair, its reduction potential (1.51 V) is more negative than that

314

of HSO5−/SO4•−, but more positive than that of HSO5−/SO5•−, which makes the

315

Mn(III)/Mn(II) redox cycle thermodynamically feasible (eqs. 3 and 7). In regard to

316

the redox pair of Fe(III)/Fe(II), its reduction potential (0.77 V) is more negative than

317

that of HSO5−/SO5•−; thus, the regeneration of Fe(II) (eq. 5) is thermodynamically

318

unfavorable.52 This was further evidenced by the obvious difference between the

319

catalytic activities of Mn3O4 and Fe3O4 (Figure 3). Specifically, our results suggest

320

that the thermodynamically favorable electron transfer from Fe(II) to Mn(III) had

321

little effect on the enhancement of BPA degradation over the as-prepared MnFeO

322

nanospheres.

323

51

Based on the standard reduction

Mn?( + e → Mn,( , @ = 1.51 V

(11)

Fe?( + e → Fe,( , @ = 0.77 V

(12)

≡Fe(II) + ≡Mn(III) → ≡Fe(III) + ≡Mn(II)

(13)

As

discussed

above,

the

redox

cycle

between

Mn(III)/Mn(II)

was

324

thermodynamically favorable, and pure Mn3O4 exhibited good performance in BPA

325

degradation. Hence, Mn was considered the main active site on the surface of the

326

MnFeO nanospheres in the activation of PMS. To further explore the role of Fe, Fe2O3

327

was introduced into the control group (Figure 5a). Similar to Fe3O4, when only Fe2O3

328

was used, BPA was hardly removed, but when Fe2O3 was combined with Mn3O4 with

329

a weight ratio of 1:1 and the total dosage of 0.1 g L–1 likewise, BPA was degraded

330

much faster than with pure Mn3O4. These results suggest that the synergetic effect still

331

existed even when Fe(II) was absent. Thus, it was Fe(III) that played a key role in the 13



332

synergism with Mn. Since the heterogeneous catalytic decomposition of peroxides

333

and most organic pollutants usually occurs on the surface of oxide particles, the

334

pre-adsorption of these reaction substrates onto the active sites, mainly through the

335

complexation effect, is especially important.35, 53-55 It was reported that the stability of

336

complexes of bivalent metal ions follows the order of Fe > Mn, irrespective of the

337

nature of the coordinated ligand or of the number of ligand molecules involved.56, 57

338

Therefore, an assumption was made that the synergetic effect between Mn and Fe

339

arose from the modification of the coordination environment of the active atoms (i.e.,

340

the Mn in the MnFeO nanospheres), which was caused by the robust coordination

341

ability of Fe (especially in the higher valence state) with PMS and other

342

oxygen-containing groups such as OH− and BPA.17, 31, 54

343

To confirm the above assumption, phosphate was introduced into the catalytic

344

systems (Figure 5b). Phosphate usually exerts a masking effect by strongly

345

coordinating with transition metals dispersed in the catalyst surface.58 As discussed

346

above, the combination of Fe2O3 and Mn3O4 substantially accelerated the degradation

347

of BPA; when an adequate amount of phosphate (10 mM) was added, however, the

348

synergistic effect completely disappeared and the BPA degradation kinetics of the

349

combination group was highly consistent with that of pure Mn3O4. This result

350

suggests that the contribution of Fe(III) in the synergism was suppressed by phosphate

351

probably through the substitution of PMS by phosphate in the coordination with

352

Fe(III). In addition, surface hydroxyl groups were considered the main factor

353

responsible for the heterogeneous catalytic activation of PMS in previous works.31, 58

354

The Fe(III) in the spinel catalyst can act as a reservoir for the hydroxyl groups and

355

donate them to a neighboring metal, which eventually facilitated the activation of

356

PMS.17, 18 In a word, the synergetic effect demonstrated in this work derives from the

357

integration of the different roles of Mn and Fe; the former acted as the main active site

358

in the catalyst surface, and the latter functioned as the main adsorption site for the

359

reaction substrates.

360

BPA Degradation Pathways. The main aromatic intermediates from BPA

361

degradation were identified as phenol, 4-isopropenylphenol, hydroquinone, resorcinol 14


Page 14 of 31

Page 15 of 31


362

and catechol by the GC-MS results (Figures S15 and S16). Based on the

363

experimental results and previous studies,52, 59-61 the possible degradation pathways of

364

BPA induced by the activation of PMS over the MnFeO nanospheres are proposed in

365

Figure S17. First, the quaternary carbon atom located in the center of BPA molecule

366

is attacked by the radicals (SO4•− and •OH) and thus phenoxyl and isopropenylphenol

367

radicals are produced through the β-scission (C-C),52 which are immediately

368

transformed

369

4-isopropenylphenol

370

4-hydroxyacetophenone as the transitional product,60-62 which is not observed in this

371

work. Meanwhile, phenol is oxidized to three types of dihydroxybenzenes through the

372

hydroxylation of the aromatic ring in the para/ortho/meta positions, and these

373

dihydroxybenzenes are further oxidized to their corresponding benzoquinones. Finally,

374

ring-opening products are formed, including muconic, maleic, oxalic, formic, acetic,

375

and malonic acids,63, 64 and finally mineralized into CO2 and H2O.

to

phenol is

and further

4-isopropenylphenol, transformed

into

respectively. hydroquinone

Second, with

376

Environmental Implications. In this work the synergistic effect between Mn and

377

Fe in their bimetallic systems for PMS activation was observed for the first time, and

378

the synergism was found to be derived from the integration of the different roles of

379

the Mn and Fe, i.e., the main catalytic active site of Mn and the main substrate

380

adsorption site of Fe. Such a synergistic effect substantially accelerated the catalytic

381

degradation of BPA. In addition to the relatively low cost and toxicity in comparison

382

to Co, Ag and Cu, the Mn-Fe bimetallic oxide has a promising potential to use PMS

383

as the oxidant for practical applications in wastewater treatment as well as in situ

384

remediation of contaminated soils and sediments, in which Mn/Fe-containing

385

minerals both exist widely. Furthermore, our findings may have important

386

implications for the rational design and effective synthesis of other Mn-Fe bimetallic

387

materials, including carbides, nitrides, etc., for sulfate radical-based advanced

388

oxidation processes. The questions of whether the synergistic effect also exists in

389

other Mn-Fe compounds and, if it does, whether its mechanism is conserved warrant

390

further investigations. Direct characterization techniques and in situ methods are also

391

needed to help elucidate the synergism between two different metal cations for PMS 15



392

activation.

393 394

AUTHOR INFORMATION

395

**Corresponding author.

396

Prof. Han-Qing Yu, Fax: +86 551 63601592; E-mail: [email protected]

397 398

Notes

399

The authors declare no competing financial interest.

400 401

ACKNOWLEDGEMENTS

402

The authors thank the National Science Foundation of China (21590812 and

403

51538011), the Collaborative Innovation Center of Suzhou Nano Science and

404

Technology of the Ministry of Education of China for supporting this work.

405 406

ASSOCIATED CONTENT

407

Supporting Information Available. The N2 adsorption-desorption isotherms (Figure

408

S1), the XRD patterns (Figure S2), the change of solution pH during the reaction

409

without any buffers (Figure S3) and with 20 mM borate buffer (Figure S4), the effect

410

of borate on BPA degradation (Figure S5), the specific-surface-area normalized

411

catalytic efficiencies (Figure S6) of the catalysts, the BPA removal efficiencies in

412

homogeneous systems (Figure S7) and in repeated batch catalytic reactions (Figure

413

S9), the TOC removal efficiency (Figure S8), the Raman (Figure S10) and FTIR

414

(Figure S11) spectra of the MnFeO samples, the SEM images of Mn3O4/Fe3O4 (Figure

415

S12), the EPR spectra in the presence of EtOH (Figure S13), the XPS spectra of Mn

416

2p and Fe 2p (Figure S14), the GC-MS chromatogram (Figure S15) and MS spectra

417

(Figure S16) of the intermediates, the proposed pathways (Figure S17) for BPA

418

degradation, the comparison between MnFeO and the previously reported

419

Mn/Fe-oxide catalysts in the catalytic Performance (Table S1), and concentrations of

420

the leaching metal ions after the reaction under various conditions (Table S2). This

421

information is available free of charge via the Internet at http://pubs.acs.org/. 16


Page 16 of 31

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422 423

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630

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Page 24 of 31

Table 1. XPS Results of the Mn 2p3/2 and Fe 2p3/2 for the MnFeO Samples Mn 2p3/2 binding energy

relative

Fe 2p3/2 binding energy

(eV)

intensity

(eV)

relative intensity MnFeO

Oct. Fe(II)/Oct. sample Mn(II)

Mn(II)/Mn(III)/

Oct.

Oct.

Tet.

Mn(IV)

Fe(II)

Fe(III)

Fe(III)

Fe(III)/Tet. Fe(III)

Mn(III) Mn(IV)

before 640.78

641.89

642.95

8:62:30

710.2

711.0

713.0

22:49:29

640.78

641.89

642.95

7:63:30

710.2

711.0

713.0

15:56:29

reaction after reaction

24


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Figure captions

Figure 1. SEM image (a), TEM image (the inset in (a)) and XRD pattern (b) of the as-synthesized MnFeO nanospheres.

Figure 2. Removal efficiency of BPA (a) and kinetic curves (b) in different reaction systems; effect of solution pH (c) and catalyst dosage (d) on BPA degradation in the PMS/MnFeO system. Reaction conditions: [BPA] = 10 mg L–1, [PMS] = 0.2 g L–1, [catalysts] = 0.1 g L–1 (for a-c), pH = 7.5 (for a, b and d) and all the solutions (except pH 4.2 in (c)) were pH buffered with 20 mM borate.

Figure 3. BPA degradation in catalytic PMS oxidation with Fe3O4-Mn3O4 mixtures as the catalysts. Reaction conditions: [catalysts] = 0.1 g L–1, [PMS] = 0.2 g L–1, [BPA] = 10 mg L–1, pH = 7.5 and buffered with 20 mM borate.

Figure 4. EPR spectra in activation of PMS under different conditions (a); effect of radical scavengers on BPA degradation in the PMS/MnFeO system (b). Reaction conditions for (a): [DMPO] = 5 mM, [PMS] = 0.02 g L–1, [MnFeO] = 0.01 g L–1, [BPA] = 1 mg L–1, and pH = 7.2; for (b): [BPA] = 10 mg L–1, [PMS] = 0.2 g L–1, [MnFeO] = 0.1 g L–1, pH = 7.5 and buffered with 20 mM borate.

Figure 5. BPA degradation in catalytic PMS oxidation with Fe2O3/Mn3O4 as the catalysts (a) and the inhibitory effect of phosphate-buffered solution (PBS) on BPA degradation in the above systems (b). Reaction conditions for (a): [catalysts] = 0.1 g L–1, [PMS] = 0.2 g L–1, [BPA] = 10 mg L–1, pH = 7.5 and buffered with 20 mM borate; for (b): [phosphate] = 10 mM, [BPA] = 10 mg L–1, [catalysts] = 0.1 g L–1, [PMS] = 0.2 g L–1, pH = 7.5 and buffered with phosphate.

25



Figure 1. SEM image (a), TEM image (the inset in (a)) and XRD pattern (b) of the as-synthesized MnFeO nanospheres.

26


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Figure 2. Removal efficiency of BPA (a) and kinetic curves (b) in different reaction systems; effect of solution pH (c) and catalyst dosage (d) on BPA degradation in the PMS/MnFeO system. Reaction conditions: [BPA] = 10 mg L–1, [PMS] = 0.2 g L–1, [catalysts] = 0.1 g L–1 (for a-c), pH = 7.5 (for a, b and d) and all the solutions (except pH 4.2 in (c)) were pH buffered with 20 mM borate.

27



Figure 3. BPA degradation in catalytic PMS oxidation with Fe3O4-Mn3O4 mixtures as the catalysts. Reaction conditions: [catalysts] = 0.1 g L–1, [PMS] = 0.2 g L–1, [BPA] = 10 mg L–1, pH = 7.5 and buffered with 20 mM borate.

28


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Figure 4. EPR spectra in activation of PMS under different conditions (a); effect of radical scavengers on BPA degradation in the PMS/MnFeO system (b). Reaction conditions for (a): [DMPO] = 5 mM, [PMS] = 0.02 g L–1, [MnFeO] = 0.01 g L–1, [BPA] = 1 mg L–1, and pH = 7.2; for (b): [BPA] = 10 mg L–1, [PMS] = 0.2 g L–1, [MnFeO] = 0.1 g L–1, pH = 7.5 and buffered with 20 mM borate.

29



Figure 5. BPA degradation in catalytic PMS oxidation with Fe2O3/Mn3O4 as the catalysts (a) and the inhibitory effect of phosphate-buffered solution (PBS) on BPA degradation in the above systems (b). Reaction conditions for (a): [catalysts] = 0.1 g L–1, [PMS] = 0.2 g L–1, [BPA] = 10 mg L–1, pH = 7.5 and buffered with 20 mM borate; for (b): [phosphate] = 10 mM, [BPA] = 10 mg L–1, [catalysts] = 0.1 g L–1, [PMS] = 0.2 g L–1, pH = 7.5 and buffered with phosphate.

30


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Table of Contents (TOC) Art

31


Degradation of Bisphenol A by Peroxymonosulfate Catalytically

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