CO2 Selectivity Enhancement and

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Alkali potassium induced HCl/CO2 selectivity enhancement and chlorination reaction inhibition for catalytic oxidation of chloroaromatics Pengfei Sun, Wanglong Wang, Xiaole Weng, Xiaoxia Dai, and Zhongbiao Wu Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.7b06023 • Publication Date (Web): 30 Apr 2018 Downloaded from http://pubs.acs.org on May 1, 2018

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Environmental Science & Technology

1

Alkali potassium induced HCl/CO2 selectivity enhancement and

2

chlorination reaction inhibition for catalytic oxidation of

3

chloroaromatics

4

Pengfei Suna, Wanglong Wanga, Xiaole Wenga,b*, Xiaoxia Daia, and Zhongbiao Wua,b

5

a. Key Laboratory of Environment Remediation and Ecological Health, Ministry of Education,

6

College of Environmental and Resource Sciences, Zhejiang University, 310058 Hangzhou, P.

7

R. China.

8

b. Zhejiang Provincial Engineering Research Centre of Industrial Boiler & Furnace Flue Gas

9

Pollution Control, 388 Yuhangtang Road, 310058 Hangzhou, P. R. China.

10

Corresponding

author:

11

[email protected].

12

Abstract

Dr.

Xiaole

Weng,

Fax/Tel:

0086-571-88982034;

E-mail:

13

Industrial combustion of chloroaromatics is likely to generate unintentional

14

biphenyls (PCBs), polychlorinated dibenzo-p-dioxins (PCDDs) and polychlorinated

15

dibenzofurans (PCDFs). This process involves a surface-mediated reaction and can be

16

accelerated in the presence of a catalyst. In the past decade, the effect of surface

17

nature of applied catalysts on the conversion of chloroaromatics to PCBs/

18

PCDD/PCDF has been well explored. However, studies on how the flue gas

19

interferent components affect such a conversion process remain insufficient. In this

20

article, a critical flue gas interferent component, alkali potassium, was investigated to

21

reveal its effect on the chloroaromatics oxidation at a typical solid acid-base catalyst,

22

MnxCe1-xO2/HZSM-5. The loading of alkali potassium was found to improve the

23

Lewis acidity of the catalyst (by increasing the amounts of surface Mn4+ after

24

calcination), which thus promoted the CO2 selectivity for catalytic chlorobenzene (CB)

25

oxidation. The KOH with a high hydrophilicity has favored the adsorption/activation

26

of H2O molecules that provided sufficient hydroxyl groups and possibly induced a

27

hydrolysis process to promote the formation of HCl. The K ion also served as a

28

potential sink for chorine ions immobilization (via forming KCl). Both of these 1 / 26

ACS Paragon Plus Environment


29

inhibited the formation of phenyl polychloride byproducts, thereby blocking the

30

conversion of CB to chlorophenol and then PCDDs/PCDFs, which potentially ensured

31

a durable operation and less secondary pollution for the catalytic chloroaromatics

32

combustion in industry.

33

Key words：：Chlorobenzene; Catalytic oxidation; alkali metal; HZSM-5; Dioxin.

34

TOC Art

35 36

1 Introduction

37

Industrial combustion of chloroaromatics emitted from solid waste incineration or

38

other thermal processes with sources of chlorine and carbon are likely to produce

39

toxic and persistent biphenyls (PCBs), polychlorinated dibenzo-p-dioxins (PCDDs),

40

and polychlorinated dibenzofurans (PCDFs) under certain reaction conditions.1-3 The

41

conversion of chloroaromatics to these persistent organic pollutants (POPs) involves a

42

surface-mediated reaction that occurs between 250 and 450 °C and can be accelerated

43

in the presence of a catalyst.4-6 As such, the catalysts and applied reaction conditions

44

for catalytic chloroaromatics combustion in industrial devices should be carefully

45

selected. Engineering technicians and/or scientists thereby require a deeper

46

understanding into the nature and mechanism of the reaction of PCBs and PCDD/F

47

precursors at potential active sites in applied catalysts. Chlorinated phenols are dominant precursors for PCDDs and PCDFs formation. 2,

48 49

6

They are usually generated by OH- nucleophilic substitution with Cl in the aromatic

50

ring of dichlorobenzene or other polychlorinated benzenes.7-9 These phenols are first

51

adsorbed onto the catalyst active site to form chlorophenolate that then either react

52

with an adjacent chlorophenol adsorbent or gas-phase chlorophenol to form 2 / 26


Page 2 of 26

Page 3 of 26


6, 10

53

PCDDs/PCDFs.

Such a reaction process could be profoundly affected by the

54

surface nature of applied catalyst, e.g., Brønsted acidity, OH- activation, and the

55

interferent components in the flue gases, e.g., fly ash, alkaline metals, NOx, and SO2.

56

In the past decade, the effect of the surface nature of applied catalyst on the formation

57

of secondary POPs has been well explored.2,

58

interferent components affect the conversion of chloroaromatics to chlorinated

59

phenols and then PCDD/PCDF remain insufficient.

11

However, studies on how the

60

In this article, we chose alkali potassium as a critical interferent component to

61

reveal its effect on the chlorobenzene (CB, a representative of chloroaromatics)

62

oxidation at a typical solid acid-base catalyst, MnxCe1-xO2/HZSM-5.12, 13 The alkalis

63

usually exist in the combustion flue gases of fossil fuels/biodiesel burning and

64

coal-fired fly ash, which can neutralize the surface acidity and reduce the redox

65

potential of applied catalysts.14,

66

potassium on the chloroaromatics oxidation process, relating to CO2 and HCl/Cl2

67

selectivity, Brønsted/Lewis acidities, and intermediate production have not been

68

thoroughly investigated. In particular, it is not yet understood how alkali potassium

69

affects the PCBs/PCDD/PCDF formation in chloroaromatics oxidation. These issues

70

require further exploration so as to provide practical guidelines for rational design of

71

industrial catalytic chloroaromatics combustion device to achieve durable operation

72

and less secondary pollution.

73

2 Materials and Methods

74

2.1 Synthesis of the catalysts

75

15

However, the interferent effect of the alkali

HZSM-5 (with an appropriate Si/Al ratio of 30

16, 17

) was supplied by Zhiyuan

76

Molecular Co., Ltd. (Shanghai, China). The Mn0.8Ce0.2O2/HZSM-5 catalyst was

77

prepared using a wet impregnation route, whereas accurately measured Mn(NO3)2,

78

Ce(NO3)3, and zeolite were mixed in ethanol with continuous stirring for 5 h. The

79

mixture was then dried at 110 °C for 10 h and calcinated at 550 °C for 5 h in static air.

80

The obtained catalyst is hereafter denoted as MCH. 3 / 26



81

Potassium-loaded

catalysts,

K(wt%)-Mn0.8Ce0.2O2/HZSM-5,

Page 4 of 26

were

also

82

synthesized using the wet impregnation route. The K(wt%) was designated as weight

83

ratio of Katom/(Katom + Mn0.8Ce0.2O2 + HZSM-5). During the syntheses, accurately

84

measured KNO3, Mn(NO3)2, Ce(NO3)3, and zeolite were mixed in ethanol with

85

continuous stirring for 5 h. The mixture was then dried at 110 °C for 10 h and

86

calcinated at 550 °C for 5 h in static air. The obtained catalysts are hereafter denoted

87

as K(1 wt%)MCH and K(10 wt%)MCH (where in the two catalysts, the weight of

88

active phase Mn0.8Ce0.2O2 was kept consistence). All of the metal salts (> 99.9%)

89

were supplied from Sinopharm Chemical Reagent Co., Ltd. and used as obtained.

90

2.2 Activity measurements and identification of intermediate products

91

The catalytic activities were measured in a fixed-bed reactor using approximately

92

1.0 g of the catalyst. The reaction feed contained 1000 ppm CB with 142 mL min-1 N2

93

and 16 mL min-1 O2 at a gas hourly space velocity (GHSV) of 10000 h-1. The reaction

94

temperature was monitored using a thermocouple loaded in the core of the catalyst

95

bed in a measuring range of 150–400 °C. All catalysts were sieved using 40-60 mesh

96

screen. The concentration of CB and CO2/CO production were analyzed on-line using

97

an Agilent 6890 gas chromatograph equipped with a flame ionization detector, an

98

electron capture detector, and nickel converting equipment.

99

The concentrations of Cl- (from HCl or Cl2) were measured using an ion

100

chromatograph instrument (Shimadzu LC-20A, Japan) equipped with a Shim-pack

101

IC-A3 adsorption column. In lab-scale measurement, the establishment of Cl. balance

102

during the catalytic oxidation of chloroaromatics is very difficult. The generated HCl

103

inclines to adsorb on the stainless steel pipe (in activity testing device) that leads to

104

very few HCl in the effluent gases. As such, measurements on the HCl or Cl2

105

production usually require an enrichment process where a 0.0125 M NaOH solution

106

was used to adsorb the HCl or Cl2 for a certain period (30 min herein). The

107

quantitative measurements (even ignoring the error) could only reveal the trend in

108

HCl or Cl2 production for each catalyst. 4 / 26


Page 5 of 26


109

The intermediate products were identified using a gas chromatography/mass

110

spectrometry (GC/MS) analyzer (the gas chromatograph was an Agilent 7890A and

111

the mass spectrometer was an Agilent 5975C). The off-gases were first captured using

112

a Tenax-GR for 30 min and then released to a thermal desorption instrument

113

(PERSEE-TP7, China) connected to a GC/MS analyzer. Quantitative measurements

114

on certain byproducts are based on a quantitative standard curve method, detail of

115

which was provided in supplementary data.

116

2.3 Catalyst characterizations

117

X-ray powder diffraction (XRD) patterns were recorded using a D/max-2500

118

diffractometer (Rigaku, Japan) with Cu Kα radiation (40 kV and 0.15418 nm). The

119

data were collected at scattering angles (2θ) from 10 to 80° with a step size of 4°. The

120

Brunauer-Emmett-Teller (BET) surface areas were determined by N2 physisorption at

121

77 K using a Micromeritics ASAP 2020 surface area and porosity analyzer. The

122

catalyst degassing pre-treatment was conducted at 150 °C for 2 h under vacuum. The

123

X-ray photoelectron spectroscopy (XPS) measurements were conducted using a

124

Thermo (America) ESCALAB 250 spectrometer with Al Kα X-ray (hν=1486.6 eV)

125

radiation as the excitation source. The charging of the catalysts was corrected by

126

setting the binding energy (BE) of adventitious carbon (C1s) to 284.6 eV.

127

The H2 temperature-programmed reduction (H2-TPR) was conducted using an

128

automatic multi-purpose adsorption instrument (TP-5079, Xianqua, Tianjin, China)

129

equipped with a custom-made thermal conductivity detector (TCD). The catalyst (50

130

mg) was first pre-treated in a purge of He at 400 °C for 1 h and naturally cooled to

131

room temperature. Then, a purge of 6 vol% H2/N2 at a flow rate of 30 mL min-1 was

132

introduced while elevating the temperature to 800 °C at a heating rate of 10 °C min-1.

133

The variation in the H2 concentration was recorded using a TCD. For O2

134

temperature-programmed desorption (O2-TPD), 0.1-g samples were preheated at

135

500 °C for 2 h under an O2 atmosphere (30 mL min-1); the samples were then slowly

136

cooled to 100 °C at a ramp of 2 °C min-1. Thereafter, the samples were swept by pure

137

He at a flow rate of 50 mL min-1 for 40 min and then heated from 100 to 900 °C. The 5 / 26



Page 6 of 26

138

signal of desorbed oxygen was recorded using a quadrupole mass spectrometer (Hiden

139

Analytical Ltd., UK).

140

Temperature-programmed surface reaction (TPSR) measurements were carried

141

out using an automatic multi-purpose adsorption instrument (TP-5079, Xianqua,

142

Tianjin, China) equipped with a quadrupole mass spectrometer (Hiden Analytical Ltd.,

143

UK). The catalyst (100 mg) was first subjected to a feed stream containing 20 vol%

144

O2/N2 and 200 ppm CB at 100 ◦C. After the CB adsorption approached equilibrium,

145

the catalyst bed was heated from 100 to 700 °C at 10 °C/min and the evolution of

146

gaseous species (including H2O, HCl, CO2, Cl2, etc.) were monitored using the MS.

147

Pyridine-adsorbed infrared (IR) spectroscopy was conducted using a Bruker

148

Fourier transform infrared (FT-IR) spectrometer equipped with a custom IR cell that

149

was connected by a vacuum adsorption apparatus. The catalyst was first heated at a

150

rate of 10 °C min-1 to 400 °C and then cooled to room temperature in a vacuum (10-3

151

Pa). Pyridine vapor was then introduced until the adsorption approached saturation.

152

The desorption process was conducted by heat-treating the adsorbed catalyst at a linear

153

heating rate of 10 °C min-1 to 450 °C. The spectra were recorded at a resolution of 4

154

cm-1. The quantitative ratio of Brønsted/Lewis was calculated as follows

155

(pyridine on B sites) = 1.88 IA(B) R2/W and C (pyridine on L sites) = 1.42 IA(L)

156

R2/W, where C is the concentration (mmol g-1 catalyst), IA (B, L) is the integrated

157

absorbance of the B or L band (cm-1), R is the radius of the catalyst disk (cm), and W

158

is the weight of the disk (mg).

159

2.4 In situ DRIFT measurements

18

: C

160

In situ diffuse reflectance infrared Fourier transform (DRIFT) spectroscopy was

161

conducted using a Nicolet 6700 FT-IR spectrometer equipped with a mercury

162

cadmium telluride detector. The DRIFT cell was equipped with CaF2 windows and

163

was fitted into a heating cartridge that allowed the catalyst to be heated to 400 °C

164

under atmospheric conditions. In each experiment, the catalyst was first pretreated in

165

a flow of He (99.99%, 100 mL min-1) at 400 °C for 1 h and then cooled to room

166

temperature. Thereafter, 100 ppm of CB and the N2 carrier gas were introduced at 6 / 26


Page 7 of 26


167

200 °C for 15 min. O2 (10 vol%) was introduced for another 15 min. The spectra

168

(average of 32 scans at 4 cm-1 resolution) were simultaneously recorded at different

169

times in each run.

170

3 Results and Discussion

171

3.1 Catalytic activity measurements

172

In activity measurements, the catalyst was subjected to a feed stream containing

173

1000 ppm CB, 10 vol% O2, and balanced N2 at a GHSV = 10000 mL gcat-1 h-1. As

174

shown in Figure 1a, the K(1wt%)MCH and MCH catalysts displayed a similar

175

catalytic activity in CB conversion, both of which revealed the T50 (i.e., temperature

176

at 50% conversion) at ~180 °C and the T90 (i.e., temperature at 90% conversion) at ~

177

240 °C. Distinct decline was observed in the K(10wt%)MCH catalyst; the T90 of the

178

catalyst dramatically increased to the temperature of ~ 450 °C. This could be caused

179

by the presence of potassium ion that excessively neutralized the acidity of HZSM-5,

180

therefore blocking the further uptake of CB over the catalyst.14, 15 After measuring the

181

CO2 production in the effluent gas, the K(1wt%)MCH catalyst was found with an

182

enhanced CO2 selectivity in respective with the MCH, whilst the K(10wt%)MCH still

183

showed a much lower CO2 selectivity. In CB-TPSR measurements (see Figure 1b),

184

the K(1wt%)MCH also displayed distinct CO2 in the temperature range of 100-300 °C,

185

where in comparison, only small humps were observed in both the MCH and

186

K(10wt%)MCH catalysts. The H2O showed a similar desorption curve to CO2, which

187

was mainly because they were both generated from the CB mineralization reaction

188

(C6H6Cl+O2→CO2+H2O+HCl).19,

189

dynamic process (i.e. the adsorption/desorption did not reach equilibrium), they still

190

reveal a consistent trend with the activity measurements for each catalyst. The

191

additional hump appeared at ~ 570 °C in the K(10wt%)MCH catalyst was mainly

192

associated with the modification of MOx (M = Ce and Mn) by K loading, which will

193

be analyzed in detail in the following section.

20

Although the CB-TPSR measurement was a

194 195 7 / 26



196 197

Figure 1 (a) Conversion of CB and selectivity of CO2 for MCH and K(wt%)MCH catalysts; (b)

198

TPSR profiles of CO2 and H2O production in MCH and K(wt%)MCH catalysts.

199

The desorption of chorine (originating from HCl and/or Cl2) from catalyst surface

200

in CB oxidation was evaluated by bubbling the outlet gas stream into a 0.0125 M

201

NaOH solution for 30 min, followed by measuring the concentrations of Cl- in

202

solution. As shown in Figure 2a, both the K-loaded MCH catalysts revealed higher

203

amounts of chorine than the MCH, suggesting that the loading of K facilitated the Cl

204

desorption from the catalyst surface. In CB-TPSR measurements; see Figure 2b, the

205

desorption amount of HCl was also much higher in the K-loaded MCH catalysts; this

206

further confirmed the ability of K in promoting the Cl desorption for catalyst surface.

207

Only little Cl2 desorption was observed in the CB-TRPS measurements, suggesting

208

that there could be no Deacon Reaction (4HCl + O2 = 2Cl2 + 2H2O) occurred in all

209

catalysts. In comparison with the input amounts of chorine, the desorbed Cl- were

210

much less in the effluent gas for both K-loaded and K-free MCH catalysts. This

211

indicated that many chorines (ignoring those adsorbed on the pipes of testing device) 8 / 26


Page 8 of 26

Page 9 of 26


212

might be residual at the catalyst surface. According to EDS-mapping (see

213

supplementary Figure S1), we did observe significant chorine species on the surface

214

of after-test K-loaded catalysts. These residual Cl species should originate from either

215

the chlorinated byproducts accumulated at surface or the generated KCl due to the

216

reaction of Cl- and K+ (the ∆H* of this reaction was only -68.0 Kcal at 298K 21).

217 218

Figure 2 (a) The amounts of absorbed chloride ions in NaOH solution for MCH and

219

K(wt%)MCH catalysts; (b) TPSR profiles of HCl and Cl2 production in MCH and K(wt%)MCH

220

catalysts.

221

3.2 Identification of gaseous byproducts

222

The gaseous byproducts in the K-loaded MCH catalysts were measured by

223

capturing the off-gas in an adsorption column (Tenax GR) at 300 °C for 30 min,

224

which were then degassed in a thermal desorption instrument and analyzed using

225

GC/MS. As shown in Figure 3 and Table 1, the K(1wt%)MCH catalyst produced

226

approximately 8 kinds of polychlorinated chain organics (labels 3, 5, 7, 10, 11, 12,

227

and 15 in Table 1) and 6 kinds of non-chlorinated chain organics (labels 1, 2, 4, 6, 8,

228

and 14). While for the K(10wt%)MCH, much less polychlorinated byproducts

229

(labeled 3, 5, and 10) were generated. In comparison with the MCH catalyst,12, 13 both 9 / 26



Page 10 of 26

230

of the K-loaded catalysts were with much less chlorinated byproducts in catalytic CB

231

oxidation. We also conducted quantitative measurements (see supplementary Figure

232

S2) on certain chlorinated byproducts in each catalyst, e.g. trichloromethane (CH3Cl),

233

trichloroethylene (C2HCl3), tetrachloromethane (C2Cl4), et al., the results of which

234

were in good agreement with the GC/MS measurements. The surface-accumulated

235

organics in the K-loaded MCH catalysts were provided in supplementary Figure S3

236

and Table S1.

237

Considering the observed significant Cl in the after-test catalyst surface and only

238

few chlorinated byproducts in effluent gases (ignoring the residual ones in pipes),

239

there is reason to believe that many chlorines could be captured by K to form KCl,

240

where the alkali K served as a sink for Cl ions immobilization. Similar phenomena

241

have been also reported in Ca modified Fe base catalysts for the catalytic

242

1,2-dichlorobenzene oxidation.22,

243

byproducts was detected in the K-loaded MCH catalysts, where in comparison, the

244

MCH

245

m-dichlorobenzene byproducts (Figure S2).12,

246

polychloride would be further converted into chlorinated phenol (as verified by in situ

247

DRIFT measurements, see Figure S4), which is the predominant precursor for PCDD

248

and PCDF formation.3, 24 The phenyl polychlorides are usually generated from the

249

electrophilic substitution of dissociated Cl. into CB at the Lewis acid sites of MClx (M

250

= metals),25 and in the K-loaded MCH catalysts, the formed KCl and the enhanced Cl

251

desorption (likely through the KOH induced H2O adsorption and subsequent H proton

252

activation that led to the hydrolysis of polychlorinated species to form HCl

253

would to some extent reduce the amounts of active Cl. at catalyst surface. This thereby

254

inhibited the formation of MnClx/CeCl4 and hence retarded the substitution reaction to

255

produce phenyl polychloride byproducts.

catalyst

was

23

however

In particular, none of the phenyl polychloride

evidenced 13

with

o-dichlorobenzene

and

This is promising as the phenyl

10 / 26


26, 27

)

Page 11 of 26


256 257

Figure 3 GC-MS analyses on the intermediate products for K(wt%)MCH catalysts

258

Table 1 Intermediate products of CB on K(wt%)MCH catalysts Label

Molecular formula

Name of Compound

1

C3H6

Propylene

2

CH2O

methanal

3

C2H3Cl

chloroethylene

4

5

C4H8

1, Butene

C4H8

2, Butene

C3H6

propylene

C2H2Cl2

dichloroethylene

Molecular structure

CH2O H

H

H

Cl

H

Cl

H

Cl

6

C3H6O

propionaldehyde

7

C3H5Cl

chloropropene

C5H10

1, Pentene

C5H10

2, Pentene

9

H2 O

water

H2O

10

CCl4

tetrachloromethane

CCl4

11

CHCl3

trichloromethane

CHCl3

o Cl

8

11 / 26



12

C2HCl3

14

3, hexene

C6H12

2, hexene

C6H12

1, hexene

C2H4O

acetaldehyde

C2Cl4

15

H

Cl

Cl

Cl

trichloroethylene

C6H12

13

Page 12 of 26

o Cl

Cl

Cl

Cl

dichloroethylene

259

3.3 Catalyst characterizations

260

3.3.1 XRD and BET surface area measurements

261

In the XRD reflection profile (see Figure 4), both of K(1wt%)MCH and

262

K(10wt%)MCH catalysts maintained a well-ordered zeolite structure. However, the

263

characteristic reflections of Mn3O4 (JCPDS 18-0803), MnO2 (JCPDS 42-1169), cubic

264

fluorite CeO2 (JCPDS 34-0394), and K2O (JCPDS 23-0493) phases were absent. This

265

could be due to the formation of ultrafine metal oxides or the high dispersion of them

266

over the catalysts.28, 29 As shown in Table 2, the K loading was found to result in a

267

consistently decrease of SBET. This could be caused by the doping of K into the

268

skeleton of HZSM-5 that blocked the micron pores of the MCH catalyst, because the

269

K(10wt%)HZSM-5 catalyst revealed significantly broadened XRD diffraction

270

patterns in the range of 5~20º and 20~50º as compared with HZSM-5 alone (see

271

supplementary Figure S5), which was reported to be caused by a typical structure

272

re-arrangement.30,

273

catalysts (see supplementary Figure S6) all displayed the characteristic curve between

274

type-IV and type-II, with hysteresis loops similar to that of type H4, indicating that

275

these catalysts were all with the slit-shaped mesopores of zeolites.32

31

The adsorption isotherms of the K-loaded MCH and MCH

12 / 26


Page 13 of 26


276 Figure 4 X-ray Powder diffraction patterns of HZSM-5, MCH, and K-loaded catalysts.

277 278

Table 2 BET surface area, Brønsted/Lewis (B/L) ratios, and XPS analyses of HZSM-5, MCH,

279

and K-loaded MCH catalysts. XPS

Brønsted

BET Catalysts

Oxygen species (%)

Mn4+/Mn

Ce3+/Ce

(m2/g)

Katom/Mnatom (wt%)

acid/Lewi

Oγ

Oβ

Oа

(%)

(%)

Theoretical

XPS analyses

s acid

HZSM-5

383.7

-

-

-

-

-

-

-

5.07

MCH

272.7

48.5

25.8

25.7

38.2

22.3

-

-

0.55

K(1wt%)MCH

244.5

58.8

32.4

8.8

48.7

19.3

0.12

0.09

0.28

K(10wt%)MCH

201.2

73.7

15.8

10.5

54.2

20.5

1.18

0.33

0.12

280

3.3.2 XPS measurements

281

In O1s XPS spectra (see Figure 5), the binding energy, BE, in the range of

282

529.4-530.0 eV is denoted to lattice oxygen species Oα; 33 the BE at approximately

283

531.9 eV is assigned to surface adsorbed oxygen and weakly bonded oxygen species

284

(Oβ, active oxygen), and the BE at approximately 532.8 eV(Oγ) corresponds to

285

oxygen-containing groups such as hydroxyl, carbonate species or adsorbed water

286

species.25,

287

Mn2p1/2 and Mn2p3/2 corresponding to a mixed-valence manganese system (Mn4+ and

288

Mn3+).35 The BE at 642.9-643.6 and 641.6-642.2 eV could be ascribed to Mn4+ and

34

The Mn2p XPS spectra generally consists of a spin-orbit doublet of

13 / 26



289

Mn3+,36 respectively. The Ce3d XPS spectra are composed of two multiplets (V and

290

U), where the peaks denoted as V, V′′, V′′′, U, U′′ and U′′′ are assigned to the 3d104f0

291

state of Ce4+, and those denoted as U′ and V′ are attributed to the 3d104f1 state of

292

Ce3+.35 The K2p XPS spectra consists of a line at 292.8 eV for K2p3/2 and a satellite

293

line at 295.5 eV for K2p1/2, both of which are attributed to K2O. 37

294 295

Figure 5 XPS analyses of O1s, Mn2p, Ce3d and K2p for the MCH and K-loaded MCH catalysts.

296

As shown in Table 2, the K(1wt%)MCH catalyst revealed much higher amounts

297

of active oxygen, Oβ (~32.4 mol%), than that of MCH (~25.8 mol%). The surficial

298

amounts of Mn4+ and Ce4+ were also enriched in this catalyst. For the K(10wt%)MCH,

299

less variety in the amounts of Oβ and Ce4+ was observed in respective with the MCH.

300

However, the Mn4+ (mol%) of the catalyst was drastically increased. The hydroxyl

301

oxygen species/adsorbed water species, Oγ (mol%), were found to distinctly increase

302

in both K-loaded catalysts. This was mainly caused by the surficial K2O that adsorbed

303

H2O from atmosphere, forming hydroxyl species. Such a H2O adsorption phenomena 14 / 26


Page 14 of 26

Page 15 of 26


304

was also observed in K2p XPS spectra, which revealed a higher BE shift in both

305

K-loaded catalysts, indicating the presence of majority KOH at catalyst surface37. The

306

KOH with a high hydrophilicity is expected to promote the H2O adsorption on

307

catalyst surface, leading to the enrichment of activated H2O molecules that either

308

bonded at metal atoms

309

These enriched H2O might consequently induce a hydrolysis process or provide

310

hydroxyl groups to promote the HCl formation in the K-loaded catalysts

311

Figure 2 for desorbed chlorine and CB-TPSR measurements).

312

3.3.3 H2-TPR measurements

38, 39

or form protonated H5O2+ dimer in the HZSM-5

40

40, 41

.

(see

313

Figure 6 illustrates the H2-TPR profile in the temperature range of 100–600 °C

314

for K-loaded MCH and MCH catalysts. The MCH catalyst displayed two H2

315

consumption humps centered at ~315 and ~368 °C, that were ascribed to the reduction

316

of Mn4+ to Mn3+ and Mn3+ to Mn2+,41, 42 respectively. The K(1wt%)MCH catalyst

317

showed similar consumption peaks at ~320 and ~372 °C, implying that the addition of

318

1 wt% K did not distinctly affect the redox potential of MCH catalyst. However, such

319

a K loading had promoted the total H2 consumption of K(1wt%)MCH catalyst, which

320

was probably due to the formation of Mn-Ce-K-O that increased the oxygen mobility

321

in the catalyst.

322

by the reduction of K2O, as by increasing the K loading to 10 wt%, the H2

323

consumption was further promoted. To verify this, a K2O/HZSM-5 catalyst was

324

prepared by solely using KNO3 and HZSM-5 as precursors. The resulted catalyst only

325

revealed a H2 reduction peak at ~638 °C (see supplementary Figure S7), which

326

verified that the increased H2 consumption in K-loaded catalysts was not caused by

327

the reduction of K2O.

43

One might argue that the enhanced H2 consumption may be caused

328

In comparison with the MCH, both of K(1wt%)MCH and K(10wt%)MCH

329

revealed an additional peak centered at ~500–535 °C. This peak was likely due to the

330

surface reduction of CeO2.44 The formation of isolated CeO2 (and possible MnOx) in

331

the K-loaded catalysts should be the result of K ions preferentially ion-exchanged

332

with the H· in HZSM-5 that hindered the ion exchange of Ce4+ and Mn3+. As such, 15 / 26



333

these ions would accumulate at the HZSM-5 surface, which could aggregate to form

334

isolated CeO2 (and MnOx) after calcination in static air. This might explain why the

335

K-loaded catalysts contained higher amounts of Mn4+ and Ce4+ (mol%) than the MCH

336

at surface (see Table 2 for XPS results).

337 338 339

Figure 6 H2-TPR profiles of the MCH and K-loaded MCH catalysts.

3.3.4 O2-TPD measurements

340

In O2-TPD profile, the oxygen desorption peaks located at < 400 °C were

341

assigned to molecular physical/chemical oxygen adsorbed on the surface, Oads; peaks

342

located between 400 and 600 °C were attributed to the oxygen adsorbed on surface

343

vacancies or the subsurface lattice oxygen, surf-Olatt, and those above 600 °C were

344

attributed to bulk lattice oxygen, bulk-Olatt.34

345 346

Figure 7 O2-TPD profiles of MCH and K-loaded MCH catalysts.

347

As shown in Figure 7, the MCH catalyst displayed three O2 desorption peaks

348

centered at approximately 412, 540, and 640 °C. The former two peaks were assigned 16 / 26


Page 16 of 26

Page 17 of 26


349

to Oads and surf-Olatt that were very active in oxidation reaction.34,

350

K(1wt%)MCH catalyst revealed a similar O2 desorption curve with the peaks centered

351

at approximately 420, 548, and 645 °C. However, the amounts of O2 desorption in the

352

catalyst were relatively higher than those in MCH catalyst. For the K(10wt%)MCH,

353

distinct O2 desorption were observed in O2-TPD measurement. The O2 desorption

354

peak at 520 °C was assigned to surf-Olatt, which should mainly originate from the

355

Mn-Ce-K-O species (see H2-TPR result in Figure 6), while the broad peak at 680 °C

356

was ascribed to the decomposition of K2O/KOH because a similar desorption peak

357

was observed in the K(10wt%)/HZSM-5 catalyst (that was prepared by using solely

358

KNO3 and HZSM-5 as precursors, see supplementary Figure S8). The presence of

359

enriched surf-Olatt in the K(10wt%)MCH catalyst should be associated with the CO2

360

and HCl humps (at ~ 550 °C) in its CB-TPSR profile (see Figure 1b and Figure 2b),

361

where at this temperature, the surf-Olatt would be involved in the CB oxidation

362

reaction that promoted the production of CO2 and HCl.

363

3.3.5 Pyridine IR measurements

364

As reported,

7, 8

45

The

the oxidation of CB over HZSM-5-based catalysts generally

365

initiated from the CB adsorption onto Brønsted acidic sites of HZSM-5 via a

366

nucleophilic substitution reaction. At this stage, the weaker C-Cl band (with respect to

367

the C-H band) in aryl halides was cleaved, which converted the CB into phenolates7, 8

368

and then benzoquinone or cyclohexanone species. Thereafter, aromatic ring cleavage

369

and deep oxidation to CO2/CO at the Lewis acidic sites of metal oxides occurred. In

370

such an oxidation process, the surface nature of Brønsted and Lewis acidity in the

371

catalysts plays a crucial role in CB oxidation efficiency.

372

Pyridine-IR measurements that were used to evaluate the change of Brønsted and

373

Lewis acidity by K loading. The bands at ~1635 and 1544 cm-1 were attributed to

374

pyridine adsorbed onto Brønsted sites, and those at ~1612 and 1455 cm-1 were due to

375

pyridine adsorbed onto Lewis sites.18, 46-48 The band at ~1490 cm-1 originated from the

376

pyridine adsorbed onto both Lewis and Brønsted sites48, 49 and the band at ~1474 cm-1 17 / 26


2

Figure 8a illustrates the


377

was ascribed to C-H bending of the aliphatic hydrocarbon bound to the oxygen atoms

378

in zeolite framework. 50, 51

379 380

Figure 8 (a) Pyridine-IR spectra analyses on the MCH and K-loaded MCH catalysts, and (b)

381

structure of the H5O2+ water dimer adsorbed over HZSM-5.

382

It was noted that the loading of 1 wt% K had distinctly promoted the Lewis

383

acidity of K(1wt%)MCH catalyst. This explained why the catalyst had a superior CO2

384

selectivity in CB oxidation (see Figure 1). Because the Lewis acidity between

385

HZSM-5 and K(1wt%)/HZSM-5 was not distinctly different (see supplementary

386

Figure S9), the incremental Lewis acidity in the K(1wt%)MCH catalyst should

387

mainly originate from the enriched Mn4+ and Ce4+ in the catalyst. As aforementioned,

388

such enrichments were caused by the inhibition of ion exchange by K loading that led

389

to the formation of enriched CeO2 and MnO2 after calcination (see Table 2 and

390

Figure 5). In addition, the presence of KOH with a high hydrophilicity might also

391

contribute to the incremental Lewis acidity (and hence CO2 selectivity) in the

392

K(1wt%)/HZSM-5 catalyst. The induced H2O adsorption could favor the formation

393

ion-pair complex (H5O2+) dimers (i.e., the protonated dimer is bound to the anionic

394

zeolite framework, see Figure 8b) in the HZSM-5, which would generate additional

395

Lewis acidity and thus increased the CO2 selectivity by forming more oxygenated

396

compounds, i.e., CxHyOz (z > 2).52, 53 In our previous works, 12, 13 we also found that

397

concentrated H2O treatment of the HZSM-5 could lead to the formation of protonated

398

H5O2+ dimers, which profoudntly increased the CO2 selecivity of MCH catalyst in CB 18 / 26


Page 18 of 26

Page 19 of 26


399

oxidation. From the Pyridine-IR profiles, we also noted that the K(1wt%)MCH

400

catalyst still retained sufficient amounts of Brønsted acidity whilst in the

401

K(10wt%)MCH, both of the Brønsted and Lewis acidities were severely reduced.

402

This is unsurprising given that the ion-exchange of excessive K would neutralize all

403

acidic sites in the K(10wt%)MCH catalyst.

404

3.4 In situ DRIFT measurements

405

To get an insight into the CB reaction process over the K-loaded catalysts, in situ

406

DRIFT measurements were further conducted while flowing CB at 200 °C for 15 min

407

and then 10 vol% O2 for another 15 min.

408

As shown in Figure 9a, after the adsorption of CB onto K(1wt%)MCH catalyst

409

at 200 °C, several bands appeared at 1030, 1130, 1170, 1250, 1320, 1370, 1461, 1530,

410

1590, and 1640 cm-1 (see Figure 9a). The bands at 1320, 1370, and 1590 cm-1

411

corresponded to -COOH from bidentate formate,

412

cm-1 were ascribed to -COOH from acetate species.55, 56 The bands at 1130 and 1250

413

cm-1 originated from the vibration of -CH2 in a different vibration model, and the band

414

at 1030 cm-1 was related to the vibration of -CH.7, 54 From the DRIFT spectra, it was

415

noted that the CB oxidation occurred upon contact with K(1wt%)MCH catalyst as the

416

bands originated from the cleavage products of aromatic ring all appeared in this

417

catalyst. The negative band at 1640 cm-1 was assigned to the CB-adsorbed on

418

Brønsted sites,18, 48 and that at 1170 cm-1 was attributed to the Lewis sites bound to

419

MnOx-CeO2.

420

adsorbed CB was continuously oxidized at the Lewis acidic sites with increasing time.

421

After introducing O2, a new band appeared at 1225 cm-1, which was attributed to the

422

C-O vibration of phenolate.7, 8 The occurrence of this band implied that the retained

423

Brønsted acidity in the K(1wt%)MCH catalyst still induced a nucleophilic substitution

424

reaction with CB, converting the CB into phenolate, and facilitating the aromatic ring

425

cleavage and deep oxidation processes. As the reaction time increased, the bands at

426

1250 (-CH2) and 1370 cm-1 (i.e., -COOH from bidentate formate) began to decrease.

427

This suggested that the aromatic ring cleavage products were gradually converted into

57

54

and the bands at 1461 and 1530

The constant consumption of these two bands implied that the

19 / 26



428

CO2/CO.

429 430

Figure 9 DRIFT spectra of (a) K(1wt%)MCH and (b) K(10wt%)MCH catalysts taken at 200 °C

431

at different times.

432

In the K(10wt%)MCH catalyst (see Figure 9b), additional bands at 1403 and

433

1510 cm-1 were observed, which originated from the vibration of benzene.58, 59 No

434

vibrational bands for phenolate were detected. The reason should be ascribed to the

435

severely neutralized Brønsted sites in this catalyst that suppressed the nucleophilic

436

substitution of CB. As such, the C-Cl band of CB had to be cleaved mainly at the

437

oxygen vacancies of MnOx-CeO2, forming phenyl radicals.58, 59 The benzene was then

438

formed via the reaction between the phenyl radical and H (provided by either Mn-OH

439

or Ce-OH according to Miran et al. by using a DFT-LDA approach), leading to the

440

appearance of characteristic benzene bands at 1403 and 1510 cm-1.19, 60 From the

441

DRIFT spectra, we deduced that the oxidation of CB over the K(10wt%)MCH

442

catalyst should be mainly through CB-benzene-CO2/HCl route, where the neutralized

443

Brønsted acidity in this catalyst had hindered the nucleophilic substitution reaction to

444

convert CB into phenolate. Because the benzene required a high activation energy for

445

aromatic ring cleavage, the K(10wt%)MCH catalyst thus revealed a poor

446

low-temperature activity than the K(1wt%)MCH catalyst in CB oxidation (see Figure

447

1).

448 449

Acknowledgement

450

This work was financially supported by the National Natural Science Foundation of 20 / 26


Page 20 of 26

Page 21 of 26


451

China (Grant No. 51478418) and the Program for Zhejiang Leading Team of S&T

452

Innovation (Grant No. 2013TD07).

453

Supporting Information Available

454

Chlorine distribution on catalyst surface, quantitative analyses of intermediate

455

compounds, DRIFT spectra for proving chlorophenol formation, etc. are in

456

supplemental section. This material is available free of charge via the Internet at

457

http://pubs.acs.org.

458

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which

are

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González-Velasco,

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for

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CO2 Selectivity Enhancement and

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