Oxidation of ethanethiol in aqueous alkaline solution by ferrate(VI): Kinetics, stoichiometry and mechanism

Accepted Manuscript Oxidation of ethanethiol in aqueous alkaline solution by ferrate(VI): Kinetics, stoichiometry and mechanism Jing Wang, Tong Zheng, Chen Cai, Yanxiang Zhang, Huiling Liu PII: DOI: Reference:

S1385-8947(18)32205-8 https://doi.org/10.1016/j.cej.2018.11.005 CEJ 20307

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Chemical Engineering Journal

Received Date: Revised Date: Accepted Date:

15 July 2018 30 October 2018 1 November 2018

Please cite this article as: J. Wang, T. Zheng, C. Cai, Y. Zhang, H. Liu, Oxidation of ethanethiol in aqueous alkaline solution by ferrate(VI): Kinetics, stoichiometry and mechanism, Chemical Engineering Journal (2018), doi: https:// doi.org/10.1016/j.cej.2018.11.005

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1

Oxidation of ethanethiol in aqueous alkaline solution by ferrate(VI): Kinetics,

2

stoichiometry and mechanism

3 Jing Wang a, Tong Zheng b,*, Chen Cai a, Yanxiang Zhang a, Huiling Liu c,*

4 5

a

6

of Technology, Harbin 150090, China

7

b

School of Environment, Harbin Institute of Technology, Harbin 150090, China

8

c

College of Environmental Science and Engineering, Tongji University, Shanghai

9

200092, China

State Key Laboratory of Urban Water Resource and Environment, Harbin Institute

10



11

E-mail: [email protected] (Tong Zheng), [email protected] (Huiling Liu).

Corresponding authors.

1

12 13

Abstract The oxidation of ethanethiol (C2H5SH), as a representative odorous pollutant, by

14

ferrate (Fe(VI)) in aqueous alkaline solution was investigated to understand its

15

kinetics, stoichiometry, and reaction pathways. The reaction kinetics and fraction of

16

different intermediates were found highly dependent on pH. The reaction behaved as a

17

5/2-order reaction in a pH range of 8.0–10.0, and a second-order reaction in a pH

18

range of 10.5–12.0. The rate constants decreased with increased pH, and the

19

speciation of Fe(VI) and C2H5SH during the reaction were further discussed. C2H5SH

20

was finally oxidized to nonvolatile and odorless sulfonic acid (C2H5SO3H) with the

21

intermediates of sulfinic acid (C2H5SO2H) and disulphide (C2H5SSC2H5). A plausible

22

mechanism involving two competing reactions for C2H5S• radicals was proposed to

23

explain the pH-dependent reaction order. In addition, kinetics and products of

24

C2H5SH degradation by Fe(VI) were compared with other oxidants (O3, chlorine,

25

ClO2, KMnO4, H2O2, and •OH). Our results provided a deep understanding of

26

C2H5SH degradation by Fe(VI) in aqueous alkaline solution.

27

Keywords: Ferrate(VI); Ethanethiol; Diethyl disulfide; Kinetics; Stoichiometry

2

28

1. Introduction

29

Odorous pollutants have gained increasing attention due to their offensive and

30

corrosive effects on the environment as well as the detrimental effect on human health.

31

Thiols, a group of typical organic sulfur compounds, are characterized by their

32

offensive odor that can be detected by human nose at very low concentrations [1-3].

33

Thiols are released from various sources such as sewage treatment works, landfill

34

sites, petroleum refining and other manufacturing processes [4-6]. Ethanethiol

35

(C2H5SH) as one of representative thiols is regarded as the smelliest substance with a

36

quite low odor threshold of ~8.7 × 10-3 ppb/v [7]. Low levels of C2H5SH in the air can

37

result in chronic harmful effects on lungs and nervous systems of human [8,9].

38

The chemical scrubbing technique has been proved to be an effective and reliable

39

process for odorous gas treatment where odorants could be absorbed to the liquid

40

phase and then oxidized rapidly by oxidants [10-12]. Ferrate(VI) (Fe(VI)) is

41

considered as a green oxidant due to its environmental benign characters and

42

integrated function of oxidation, disinfection and coagulation [13,14]. It is also a

43

selective oxidant exhibiting high reactivity with electron-rich groups of target

44

compounds including organic sulfur compounds, amines and phenols [15,16]. Fe(VI)

45

has shown potential advantages in sludge odor control for sewage treatment works

46

(e.g., thiosemicarbazide , thiourea dioxide and methyl mercaptan) [17]. Furthermore,

47

there is evidence that gaseous methyl mercaptan can be effectively degraded by the

48

electrochemically generated Fe(VI) in NaOH solution [18,19]. The rate constants of 3

49

thiols oxidation by Fe(VI) were reported in the range of ~104 M-1 s-1 at pH = 9 [20].

50

Therefore, C2H5SH with low molecular weight and weak acidity (pKa = 10.5 [21])

51

might be absorbed readily into the alkaline Fe(VI) solution, and then be oxidized with

52

an extremely fast reaction rate.

53

Significant progress has been made in exploring the kinetics and mechanism of

54

the reaction between Fe(VI) and organic pollutants in neutral and alkaline conditions.

55

Ordinarily, the role of pH in kinetics of the reaction is embodied in the variation of

56

rate constants with pH which could generally be quantitatively modeled [16,22,23].

57

However, the possible effect of pH on intermediates and products of the reaction

58

requires investigation as well. Reaction mechanisms involving one and two-electron

59

transfer as well as oxygen transfer steps have been proposed in the oxidation by

60

Fe(VI), depending on the nature of organic pollutants [20,24]. Nevertheless, more

61

information on radical formation and the intermediate species of Fe(VI) is necessary

62

for a deeper understanding of the reactions, which may assist with distinguishing

63

electron transfer and oxygen transfer steps [15,25]. Studies have been carried out on

64

the mechanism of thiols oxidation by Fe(VI), but individual thiols differ markedly in

65

their behavior toward Fe(VI) [20]. Methyl mercaptan oxidation by Fe(VI) might go

66

through a one-electron transfer step to produce Fe(V) and thiyl radical in the first step.

67

The final products were Fe(III) and SO42- with the cleavage of C-S bond [17]. The 2-

68

mercaptoethanesulfonic acid could be oxidized to sulfonic acid through two-electron

69

transfer steps, and Fe(VI) was finally reduced to Fe(III) with the initial product of 4

70

Fe(II) [26]. Cysteine oxidation occurred through two-electron transfer steps as well

71

producing cysteine sulfinic acid, but Fe(II) was detected as the only final product of

72

Fe(VI) [20,27]. Therefore, more studies are needed to further understand the oxidation

73

of thiols by Fe(VI). However, we are not aware of any studies that investigate on the

74

mechanism of C2H5SH oxidation by Fe(VI) in aqueous alkaline solution. Moreover,

75

dependence of reaction order and intermediates fraction on pH have not been reported

76

in such a reaction, which are not very common in the Fe(VI) reactions.

77

Hence, the objectives of this study are to (i) investigate the effect of pH on

78

reaction kinetics of C2H5SH with Fe(VI) in aqueous alkaline solution; (ii) determine

79

the reaction stoichiometry and identify the products of C2H5SH; (iii) propose a

80

plausible mechanism to explain the changes in reaction orders at different pH; and (iv)

81

compare the kinetics and products of C2H5SH degradation by Fe(VI) with other

82

known oxidants.

83

2. Materials and methods

84

2.1. Chemicals and reagents

85

C2H5SH (GC grade, ≥ 99.5% purity), diethyl disulfide (C2H5SSC2H5, GC grade, >

86

99.0% purity) and sodium ethanesulfinate (C2H5SO2Na, 93% purity) were purchased

87

from Aladdin Chemical Inc. (Shanghai, China). Sodium ethanethiolate (C2H5SNa, 90%

88

purity) and ethanesulfonic acid (C2H5SO3H, 95% purity) were purchased from Sigma

89

Aldrich Chemical Inc. Potassium ferrate (K2FeO4) was synthesized according to the

90

wet method described by Li et al. [28]. Briefly, K2FeO4 was resulted from the 5

91

oxidation of ferric nitrate (Fe(NO3)3·9H2O) by potassium hypochlorite (KClO) in

92

strong alkaline solution, which was then further purified to obtain a higher purity.

93

Solid K2FeO4 was then dried and stored in a vacuum desiccators. The purity of solid

94

K2FeO4 was determined to be 95.6% by the chromite analytical method [29]. Further

95

details on the procedures of synthesis and purity analysis are provided in

96

supplementary material (Texts S1), as well as the characterization of K2FeO4 by X-ray

97

powder diffraction (XRD) (Text S2). Other chemicals used in the experiments were of

98

analytical grade and used without further purification. All aqueous solutions were

99

prepared with ultrapure water obtained from a Milli-Q water purification system (18

100

MΩ cm, Millipore, USA), and the pH was adjusted by adding either phosphoric acid

101

or NaOH solution.

102

The aqueous solution of Fe(VI) was freshly prepared by dissolving solid K2FeO4

103

in pure water or buffer solution, and subsequently filtrated through a 0.45μm

104

polyethersulfone (PES) syringe filter (Tianjin NAVIGATOR, China). Fe(VI)

105

concentration was determined at 510 nm with ε510 nm = 1150 M-1 cm-1 using a UV-Vis

106

spectrophotometer [30]. Aqueous solutions of C2H5SH were freshly prepared by

107

either dissolving solid C2H5SNa or adding C2H5SH to pure water just before use. The

108

stock solution of C2H5SSC2H5 for kinetic studies was prepared according to the

109

following steps. 0.5 mL C2H5SSC2H5 was added into 1 mL methanol, and then 0.25

110

mL tween 80 was added into the mixture to increase the aqueous solubility of

111

C2H5SSC2H5. After that, water was slowly added and the solution volume was 6

112

adjusted to 50 mL. C2H5SSC2H5 solution for products studies was prepared in

113

methanol.

114

2.2. Kinetics studies

115

Kinetics studies on the reaction of Fe(VI) with substrates (C2H5SH and

116

C2H5SSC2H5) were investigated using a stopped-flow spectrophotometer (SX 20)

117

from Applied Photophysics Ltd, UK at 25 ± 1 °C. The Fe(VI) solutions were prepared

118

in 5 mM Na2HPO4/1 mM Na2B4O7·10H2O buffer solutions (pH = 9.2). The aqueous

119

solutions of substrates were prepared in buffer solutions with a pH range of 7.0‒12.0.

120

20 mM phosphate buffer solution was used for pH = 7.0‒8.0. 10 mM borate buffer

121

solution was used for pH = 8.5‒12.0 together with 20 mM Na2HPO4. Phosphate was

122

used to control pH as well as to avoid the precipitation of Fe(III). The Fe(III)

123

precipitation not only interferes with the optical monitoring of the reaction but also

124

causes damage to the instruments. The experiments were performed under the pseudo-

125

order condition with substrates in at least 10-fold excess. The Fe(VI) concentrations

126

ranged from 0.09 to 0.14 mM. Fe(VI) solutions were mixed with equal volumes of

127

substrate solutions at the studied pH. Reaction rate constants were determined by

128

monitoring the decay of Fe(VI) absorbance as a function of time. Five experimental

129

runs were performed for each concentration of substrates to ensure the reliability of

130

the data. The pseudo-order rate constants reported in this study represented the

131

average values with the relative standard deviations less than 5%. In the control

132

experiments for Fe(VI) reaction with C2H5SH, the decrease in the concentration of 7

133

Fe(VI) in the given reaction time was less than 5% of the initial Fe(VI) concentration

134

at each studied pH. Therefore, the self-decomposition of Fe(VI) was ignored in

135

evaluating the rate constants. The pseudo first-order rate constants for the reaction

136

between Fe(VI) and C2H5SSC2H5 were corrected for the Fe(VI) decreases in the

137

presence of tween 80 (< 0.1%, v/v) at each studied pH. The modeling of the kinetic

138

data was performed with the software of Matlab R2016b.

139

2.3. Stoichiometry and products studies

140

Stoichiometric experiments were carried out at pH = 9.2 and 12.0 in 10 mM

141

borate/20 mM Na2HPO4 buffer solutions by mixing equal volumes of Fe(VI) and

142

C2H5SH solutions. Buffer solutions were used to maintain the pH of the reaction

143

mixtures. C2H5SH solutions were prepared at a constant concentration of 2.0 mM,

144

while Fe(VI) solutions ranged from 0 to 6.2 mM with different molar ratios of Fe(VI)

145

to C2H5SH in the reaction mixture. In the experiments without buffer, aqueous

146

solutions of Fe(VI) and C2H5SH were prepared in pure water at pH = 9.2 and 12.0 and

147

mixed immediately after preparation. Each reaction was conducted for 15 min at 20 ±

148

1 °C and the pH of the mixed solution was measured. The concentrations of C2H5SH

149

were determined according to the Ellman’s reagent method described in detail by

150

Riener et al. [31]. The concentrations of sulfonic ion (C2H5SO3-) and sulfinic ion

151

(C2H5SO2-) as the products of C2H5SH degradation were quantified by Dionex ICS-

152

2100 ion chromatography (IC) with an IonPac AS11-HC anion column (4 × 250 mm).

153

The final oxidation state of iron was determined using o-phenanthroline and 8

154

thiocyanate.

155

To identify other possible intermediates, the reaction solution was first extracted

156

with methyl tert-butyl ether (MTBE, GC grade, ≥ 99.9%), and then the extractant was

157

analyzed by gas chromatography/mass spectrometry (GC/MS) (Agilent 6890GC-

158

5975MSD) with a HP-5MS column (30 m × 0.25 mm). The injection mode was

159

splitless injection and a flow rate of the carrier gas (Helium) was 1.0 mL min-1. The

160

temperature of GC oven was set at 30 °C for 10 min, and then ramped up to 315 °C at

161

10 °C min-1. Electron ionization mode with an ionization energy of 70 eV was used in

162

mass spectrometer, and the scanned mass ranged from 20 to 400 m z-1 in a full scan

163

mode. The possible intermediates were analyzed and matched with the NIST standard

164

reference database. The concentration of C2H5SSC2H5 in the extractant was quantified

165

by GC/MS with selected ion monitor, and the extraction recovery was determined to

166

be ~95%.

167

3. Results and Discussion

168

3.1. Kinetics

169

The reaction rate of Fe(VI) with substrate is expressed and calculated by Eq. (1):

170

- d[Fe(VI)]/dt = k [Fe(VI)]m[S]n

171

where [Fe(VI)] and [S] are the concentrations of Fe(VI) and substrates (C2H5SH and

172

C2H5SSC2H5), respectively, M; k represents the apparent rate constant, M1-m-n s-1;

173

while m and n are the orders in the concentrations of Fe(VI) and substrate,

174

respectively.

(1)

9

175 176

The experiments were carried out with substrates in large excess. Eq. (1) could be rewritten as Eq. (2):

177

- d[Fe(VI)]/dt = k' [Fe(VI)]m

178

(2)

179

where k' = k [S]n, M1-m s-1.

180

The decreases in the concentrations of Fe(VI) with time at different pH are

181

shown in Fig. S3. In the pH range of 8.0–10.0, plots of ([Fe(VI)]0/[Fe(VI)]t)1/2 versus

182

time at different C2H5SH concentrations were found to be almost linear with positive

183

slopes and intercepts (intercept = 1) (R2 > 0.98, Fig. S4(a)). [Fe(VI)]0 and [Fe(VI)]t

184

represent the concentrations of Fe(VI) (M) at time zero and any time t, respectively.

185

The results indicated that the reaction was 3/2-order with respect to Fe(VI), i.e., m =

186

3/2. This was different from the common results that the reaction between Fe(VI) and

187

organic sulfur compounds showed a first-order dependence on each reactant [20,25].

188

It might be the first time that a fraction order with respect to Fe(VI) was obtained in

189

the reaction of Fe(VI) with thiols. However, in the pH range of 10.5–12.0, the

190

decreased absorbance of Fe(VI) as a function of time fitted well to single exponential

191

model, indicating that the reaction showed the first-order dependence on Fe(VI), i.e.,

192

m = 1 (R2 > 0.99, Fig. S4(b)). The k' values which were obtained from the linear

193

regression at different C2H5SH concentrations increased linearly with the C2H5SH

194

concentrations in both the pH ranges (R2 > 0.98, Fig. S5). This indicated that the

195

reaction was first-order in C2H5SH, i.e., n = 1. Moreover, the slopes n obtained from 10

196

the linear regression of logk' versus log[C2H5SH] also confirmed the first-order

197

dependence on C2H5SH (Table S1). Thus, two rate laws for the reaction of Fe(VI)

198

with C2H5SH could be expresses as Eqs. (3) and (4), respectively:

199

- d[Fe(VI)]/dt = k [Fe(VI)]3/2[C2H5SH]

200

- d[Fe(VI)]/dt = k [Fe(VI)][C2H5SH]

201

pH = 8.0–10.0

(3)

pH = 10.5–12.0

(4)

The values of k decrease nonlinearly from (5.29 ± 0.30) × 106 M-3/2 s-1 at pH =

202

8.0 to (1.58 ± 0.06) × 106 M-3/2 s-1 at pH = 10.0 (Fig. 1 and Table S1). This pH

203

dependence could be explained by the acid-base equilibrium of mono-protonated

204

Fe(VI) (HFeO4-) and C2H5SH (Eqs. (5) and (6)).

205

HFeO4 - ↔ H + + FeO42 -

pKa,

HFeO4 -

= 7.23 [32]

206 207

(5) C2H5SH ↔ H + + C2H5S -

pKa,

C2H5SH

= 10.5 [21]

208 209

(6) In the pH range of 8.0–10.0, two Fe(VI) species and two C2H5SH species could

210

react with each other (HFeO4- + C2H5SH, HFeO4- + C2H5S-, FeO42- + C2H5SH and

211

FeO42- + C2H5S-). The kinetic model described by Eqs. (7) and (8) was thus used to

212

stimulate the values of k at pH = 8.0–10.0. - d[Fe(VI)] dt = k [Fe(VI)]3/2 tot [C2H5SH]tot

213 = 214

∑kα ij

3/2 i

βj [Fe(VI)]3/2 tot [C2H5SH]tot

i = 1, 2 j = 1, 2

(7)

11

k

∑kα

= 215

ij

3/2 i

βj

i = 1, 2 j = 1, 2

(8) 216

where [Fe(VI)]tot = [HFeO4-] + [FeO42-], M; [C2H5SH]tot = [C2H5SH] + [C2H5S-], M;

217

αi and βj are the species fraction for Fe(VI) and C2H5SH, respectively; i and j are each

218

of the two Fe(VI) species and two C2H5SH species, respectively; and kij is the species-

219

species rate constant for the reaction of Fe(VI) species with C2H5SH species, M-3/2 s-1.

Measured k Model k

7

kapp (M-3/2 s-1)

10

106 5

10

k21FeO

4

3/2 2-

βC H SH 2

k12HFeO

5

3/2 -

4

k11HFeO

3/2 -

4

4

βC H S 2

-

5

βC H SH 2

5

10

103

7.5

8.0

8.5

9.0

9.5

10.0

pH

220 221

Fig. 1. Apparent rate constants k for the reaction of Fe(VI) with C2H5SH as a function

222

of pH at 25 ± 1 °C.

223

To further simplify the kinetic model, the reaction of FeO42- with C2H5S- was

224

neglected in model fitting due to its negligible contribution to the overall reaction.

225

The critical evaluation of the exclusion of FeO42- reaction with C2H5S- in the model

226

fitting is provided in supplementary material (Text S3). The values of the species-

227

species rate constants were estimated by the least-squares nonlinear regression of the 12

228

measured k values. The results fitted reasonably well with the experimental data (R2

229

＞ 0.99, solid line in Fig. 1). The estimated species-species rate constants were k11 =

230

(2.11 ± 0.23) × 107 M-3/2 s-1 for the reaction of HFeO4- with C2H5SH, k12 = (1.56 ±

231

0.11) × 1010 M-3/2 s-1 for the reaction of HFeO4- with C2H5S- and k21 = (1.77 ± 0.07) ×

232

106 M-3/2 s-1 for the reactions of FeO42- with C2H5SH, respectively. HFeO4- reacted

233

faster with protonated C2H5SH than FeO42-. Density functional theory calculations

234

showed that HFeO4- had higher oxidizing power than FeO42- owing to its lower

235

LUMO energy level in aqueous solution [33]. The fraction of HFeO4- decreased with

236

increased pH, which might be responsible for the decrease in the rate constants

237

between Fe(VI) and C2H5SH with the increased pH. Moreover, HFeO4- reacted with

238

C2H5S- approximately three orders of magnitude faster than with protonated C2H5SH.

239

This was in agreement with the previous investigations that thiolates, which were

240

orders of magnitudes more reactive as a nucleophile than the corresponding thiols,

241

were much more prone to be oxidized than protonated thiols [34]. Contributions of the

242

specific reactions to the apparent rate constants are shown in Fig. 1 (Dashed lines).

243

While the reaction of HFeO4- with C2H5SH dominated at neutral pH, both the

244

reactions of HFeO4- with C2H5S- and FeO42- with C2H5SH might dominate at pH =

245

7.7–8.6 and pH = 8.6–10.0, respectively.

246

As shown in Table S1, the values of k at pH = 10.5–12.0 were also pH dependent.

247

It slightly decreased from (1.30 ± 0.02) × 104 M-1 s-1 at pH = 10.5 to (0.93 ± 0.03) ×

248

104 M-1 s-1 at pH = 12.0, where FeO42- dominated the reaction. The reaction rates at 13

249

pH = 10.5–12.0 are influenced by two factors of (1) the fraction of C2H5S- with a

250

higher reaction rates than C2H5SH and (2) the oxidizing power of FeO42-. The

251

calculations of redox potential showed the decrease in electrode potential of FeO42-

252

/Fe(OH)3 couple with increased pH in alkaline conditions, which indicated that the

253

oxidizing power of FeO42- decreased as well [35]. Therefore, the decrease of the k

254

value as the increased pH at pH = 10.5–12.0 was attributed to the decreased oxidizing

255

power of FeO42- [25,35].

kapp (M-1 s-1)

40

Measured k Model k

30 20 10 0 7.0

7.5

8.0

8.5

9.0

9.5

10.0

10.5

pH

256 257

Fig. 2. Apparent rate constants k for the reaction of Fe(VI) with C2H5SSC2H5 as a

258

function of pH at 25 ± 1 °C.

259

The kinetics for the oxidation of C2H5SSC2H5, an intermediate of C2H5SH

260

degradation by Fe(VI), was investigated as well. The results indicated first-order

261

dependence on both of Fe(VI) and C2H5SSC2H5 (Fig. S6 and Eq. (9)). The values of k

262

for C2H5SSC2H5 which significantly decrease with the increased pH show similar pH

263

dependence as for C2H5SH (Fig. 2). The analysis of k values was performed using the 14

264

kinetic model of Eqs. (7) and (8), in which the reaction order in Fe(VI) was first-order

265

instead. Two reactions between two Fe(VI) species and C2H5SSC2H5 were considered

266

in model fitting (Text S3). The estimated species-species rate constants for the

267

reaction of HFeO4- with C2H5SSC2H5 and FeO42- with C2H5SSC2H5 were 58.47 (±

268

0.51) M-1 s-1 and 0.55 (± 0.16) M-1 s-1 (R2 > 0.99), respectively. It has been reported

269

that the reaction rate of organic sulfur compounds with Fe(VI) was controlled by the

270

degree of nucleophilicity of the sulfur atom [20,25]. The decreased nucleophilicity of

271

sulfur in C2H5SSC2H5 than C2H5SH was thus responsible for the much lower reaction

272

rate of C2H5SSC2H5. The rate constant for C2H5SSC2H5 with H2O2 was reported to be

273

(1.1 ± 0.4) × 10-4 M-1 s-1 at pH = 9.0 [21]. Hypochlorous acid (HOCl) was proved to

274

have a high reactivity with disulfides, and the rate constant for 3,3'-dithiodipropionic

275

acid with HOCl was (1.6 ± 0.6) × 105 M-1 s-1 at pH = 7.2–7.4 [36,37]. Overall, the

276

reaction rate of Fe(VI) with C2H5SSC2H5 was orders of magnitude higher than the rate

277

for H2O2 in alkaline solution, while probably much lower than the rate for HOCl in

278

neutral solution.

279

- d[Fe(VI)]/dt = k [Fe(VI)][C2H5SSC2H5]

280 281 282

(9) 3.2. Stoichiometry and Products The kinetic results indicated that the reaction of Fe(VI) with C2H5SH showed

283

overall 5/2-order behavior at pH = 8.0–10.0 but second-order behavior at pH = 10.5–

284

12.0. To investigate the possible impact of pH on the intermediates and products of 15

285

C2H5SH oxidation, the stoichiometric experiments were carried out at pH 9.2 and 12.0

286

which were in the pH range of 8.0–10.0 and 10.5–12.0, respectively. The degradation

287

of C2H5SH and its oxidized products by Fe(VI) at pH 9.2 and 12.0 in buffer solutions

288

are shown in Figs. 3 and 4. An increase in the amount of Fe(VI) resulted in a

289

nonlinear reduction of the C2H5SH concentration at both pH = 9.2 and 12.0. The

290

degradation of 1 mol C2H5SH needed ~1.2 mol Fe(VI) at pH = 9.2 but ~1.4 mol Fe(VI)

291

at pH = 12.0, although the final oxidation of C2H5SH had not been achieved. However,

292

2.2 mol Fe(VI) was found to be sufficient to completely convert 1 mol C2H5SH to its

293

final product of C2H5SO3H, but it required more than 2.4 mol Fe(VI) at pH = 12.0.

294

These results also indicated that Fe(VI) was not only consumed by C2H5SH but also

295

by other intermediates from its initial degradation reactions.

296

In the oxidation process, the formation of oxidized products varied with pH and

297

the amount of Fe(VI). At pH = 9.2, C2H5SO3H was gradually formed as a final

298

oxidized product (Fig. S7(a)), and its formation was slow at the beginning of the

299

reaction and then speeded up as the concentration of Fe(VI) raised up to 4.4 mM. In

300

the meantime, two relatively stable intermediates were formed before C2H5SO3H

301

formation, and they were determined to be C2H5SO2H and C2H5SSC2H5 (m/z = 122)

302

by IC and GC/MS, respectively (Fig. S8). As shown in Fig. 3, the amount of

303

C2H5SO2H formation significantly increased with the increased Fe(VI) concentration

304

of up to 2.4 mM, and then decreased with a further increased concentration of Fe(VI).

305

A same trend was also observed for the C2H5SSC2H5 formation. 16

306

To determine if C2H5SO3H could be further oxidized by Fe(VI) to its inorganic

307

form of SO42- at pH = 9.2 and ambient temperature, experiments were conducted by

308

mixing excess Fe(VI) with C2H5SO3H in buffer solutions for more than 12 hours. The

309

initial concentration of C2H5SO3H in the reaction mixture was kept constant at 0.1

310

mM and the concentrations of Fe(VI) were 0.6 and 1.0 mM, respectively. At both

311

conditions, the degradation of C2H5SO3H was found to be less than 5% and no SO42-

312

was obviously detected in the reaction mixture detected by IC. Therefore, C2H5SO3H

313

was probably the final oxidation product of C2H5SH at pH = 9.2 and ambient

314

temperature.

Species (mM)

2.0 pH = 9.2

1.5

1.0

(1)

C2H5SO2H

(2)

C2H5SSC2H5 (3) C2H5SO3H

0.0 0.0

(4)

S - Sum(1 + 2 + 3 + 4)

0.5

315

C2H5SH

1.0

2.0

3.0

4.0

5.0

6.0

7.0

Fe(VI) (mM)

316

Fig. 3. Degradation of C2H5SH and its oxidized products by Fe(VI) at pH = 9.2 in

317

buffer solution. Experimental conditions: [C2H5SH]0 = 2.0 mM, [Fe(VI)] = 0–6.2 mM.

318

At pH = 12.0, C2H5SO2H is a major product equivalent to almost 95% of the

319

consumed sulfur when the concentration of Fe(VI) is 2.8 mM, while C2H5SO3H is a

320

minor product equivalent to only 5% (Figs. 4 and S7(b)). The formation of C2H5SO2H 17

321

increased steadily with an increased Fe(VI) concentration of up to 2.8 mM, and then

322

demonstrated a slight decrease with a further increased Fe(VI). Furthermore,

323

C2H5SSC2H5 was initially formed and then gradually disappeared as the amount of

324

Fe(VI) increased. At both pH values, the sulfur mass balance in the reaction mixture

325

was maintained.

326

Experiments were also carried out at pH = 9.2 and 12.0 without buffer, in which

327

equal volumes of 5.0 mM Fe(VI) and 2.0 mM C2H5SH solutions were rapidly mixed.

328

At pH = 9.2, C2H5SO3H was the only products of C2H5SH. At pH = 12.0, C2H5SO2H

329

was the main product and C2H5SO3H was minor, which were similar with the results

330

obtained in buffer solutions. This indicated that the type of final products for C2H5SH

331

oxidation by Fe(VI) was not influenced by the existence of buffer. The final product

332

of Fe(VI) reduction was detected as Fe(III), in which Fe(OH)3 precipitated after the

333

reaction. Besides, the pH of the reaction mixture without buffer increased from the

334

initial value of 9.2 to 11.0 and from 12.0 to 12.1. 2.0

1.5

1.0

C2H5SH

(1)

1.5

C2H5SO2H

(2)

1.0

C2H5SSC2H5 (3)

0.5

C2H5SO3H

2.0

C2H5SSC2H5 (3)

0.0 0.0

0.5

0.0 0.0

335

Species (×10-2 mM)

Species (mM)

pH = 12.0

1.0

2.0

3.0

4.0

Fe(VI) (mM)

1.0

2.0

3.0

4.0

Fe(VI) (mM)

18

5.0

(4)

S - Sum(1 + 2 + 3 + 4)

336

Fig. 4. Degradation of C2H5SH and its oxidized products by Fe(VI) at pH = 12.0 in

337

buffer solution. Experimental conditions: [C2H5SH]0 = 2.0 mM, [Fe(VI)] = 0–4.7 mM.

338

3.3. Plausible Mechanism

339

The plausible pathways of C2H5SH oxidation by Fe(VI) in aqueous alkaline

340

solution are shown in Fig. 6. Fe(VI) can oxidize thiols through a one-electron transfer

341

or two-electron transfer step to result in different products [20]. Reactions underwent

342

two-electron transfer steps would always yield Fe(II) as a reduced product from

343

Fe(VI), which were contrary to our experimental results. Furthermore, the observed

344

stoichiometric ratios of Fe(VI) to C2H5SH for complete oxidation of C2H5SH to

345

C2H5SO2H and C2H5SO3H were approximately 2.2 and 1.4, respectively. The results

346

were similar to the conclusion that the stoichiometric coefficient of Fe(VI) to organic

347

sulfur compound was 0.67 for a one oxygen-atom transfer when Fe(III) was the

348

reduced product [25]. Therefore, the initial step of C2H5SH oxidation by Fe(VI) might

349

take place to give C2H5S• radical and Fe(V) through a one-electron transfer step (Eq.

350

(10)). Moreover, the detection of C2H5SSC2H5 in the products of C2H5SH oxidation

351

provided additional support for the yield of C2H5S• radical due to the rapidly

352

recombination of alkylthiyl radicals (2kt = ~ 109 M−1 s−1 at pH = 5‒6 [38]). Then,

353

further oxidation of the C2H5S• radicals take place through two possible reactions

354

(Eqs. (11) and (12)), which are further reactions with Fe(VI) to form ethyl sulfenic

355

acid (C2H5SOH) and the dimerization of C2H5S• radicals. This two reactions are

19

356

assumed as the rate-controlling steps leading to the consumption of Fe(VI) and

357

oxidation of C2H5SH.

358

Fe(VI) + C2H5SH → C2H5S• + Fe(V)

359

k1

(10) k2

360

2C2H5S• → C2H5SSC2H5

(11)

361

C2H5S• + Fe(VI) + H2O → Fe(V) + C2H5SOH

k3

362

(12)

363

By applying the steady-state assumption to the concentration of C2H5S• radical,

364

Eq. (13) will be obtained.

365

k1[Fe(VI)][C2H5SH] = k2[C2H5S•]2 + k3[C2H5S•][Fe(VI)][H2O] (13)

366

where k1, k2 and k3 are the rate constants for the reactions of Eqs. (10)‒(12),

367

respectively, M-1 s-1; [Fe(VI)], [C2H5SH], [C2H5S•] and [H2O] represent the

368

concentrations for each of the reactants involved in the reactions of Eqs. (10)‒(12), M.

369 370

371

372 373 374

If k2[C2H5S•]2 ≫ k3[C2H5S•][Fe(VI)][H2O] is a reasonable assumption at pH = 8.0–10.0 when C2H5SH is in excess, Eq. (14) will be obtained. - d[Fe(VI)] dt = k3[C2H5S•][Fe(VI)][H2O] = k4[Fe(VI)]3/2[C2H5SH]1/2 (14) k1 1/2

where k4 = k3[H2O](k ) , M-1 s-1. 2 When [C2H5SH] ≫ [Fe(VI)], for any given concentration of Fe(VI), Eq. (14) reduces to Eq. (15), and this equation is identical with the experimental kinetic 20

375

376

expression for Fe(VI) at pH = 8.0–10.0. - d[Fe(VI)] dt = k5[Fe(VI)]3/2 (15)

377 378

k1 1/2

where k5 = k3[H2O]([C2H5SH]k ) , M-1/2 s-1. 2 If k2[C2H5S•]2 ≪ k3[C2H5S•][Fe(VI)][H2O] is a reasonable assumption at pH

379

= 10.5–12.0, Eq. (16) will be obtained, which is identical with the experimental

380

kinetic expression at pH = 10.5–12.0.

381

- d[Fe(VI)] dt = k3[C2H5S•][Fe(VI)][H2O] = k1[Fe(VI)][C2H5SH] (16)

382

As shown in Figs. 3 and 4, the amount of C2H5SSC2H5 formed at pH = 12.0 was

383

much less than the results at pH = 9.2 as the amount of Fe(VI) increased. It is

384

supposed that the increased pH is disadvantageous for the reaction of Eq. (11), which

385

is in accordance with the assumptions above. The reaction order with respect to

386

C2H5SH in Eq. (14) is inconsistent with our kinetic results that the reaction shows

387

approximate first-order dependence on C2H5SH at pH = 8.0–10.0. However, the

388

results as shown in Table S1 reveal that a reaction order with respect to C2H5SH tends

389

to decrease with a decreased pH value. This is in accordance with the results

390

speculated in Eqs. (14) and (16). At pH = 10.0–10.5, it is likely that the reaction is

391

mixed first and 3/2 order with respect to Fe(VI). No attempt has been made in this

21

392

study to analyze the data obtained in this intermediate pH region, and it needs a

393

further investigation.

394

It has been reported that sulfenic acid does not permit isolation except in very

395

special cases due to its very high reactivity [39], and therefore it would further be

396

oxidized rapidly to form C2H5SO2H. Fe(V) is approximately orders of magnitude

397

more reactive with compounds than Fe(VI) [40]. Hence, it is possible that once Fe(V)

398

formed, it could then be consumed immediately .

399

C2H5SSC2H5, still an odorous pollutant with its odor threshold value ranged from

400

0.05 to 16.5 ppb [2], can be degraded by Fe(VI) to C2H5SO2H. Reaction of Eq. (18)

401

was studied at pH = 9.2 in buffer solution by mixing Fe(VI) with C2H5SSC2H5 at

402

different molar ratios. The initial concentration of C2H5SSC2H5 in the reaction

403

mixture was kept constant at 0.4 mM and the concentrations of Fe(VI) ranged from 0

404

to 0.4 mM. Both of C2H5SO2H and C2H5SO3H were detected as the products of

405

C2H5SSC2H5. With the increased amount of Fe(VI), C2H5SO2H was initially formed

406

and gradually increased, and then C2H5SO3H was formed (Fig. 5). The oxidation of

407

C2H5SSC2H5 might occur in two possible pathways [21,36]. In the first pathway,

408

Fe(VI) oxidizes C2H5SSC2H5 by attacking on its S-S bond to yield C2H5SOH with a

409

cleavage of S-S bond, and then C2H5SOH further reacts with Fe(VI) rapidly to yield

410

C2H5SO2H. In the second pathway, Fe(VI) initially oxidizes C2H5SSC2H5 to diethyl

411

thiolsulfinate (C2H5SOSC2H5) or diethyl thiolsulfonate (C2H5SO2SC2H5) that is

412

analogous to the oxidation of cystine by Fe(VI) to thiosulfonate [27], and then further 22

413

attacks on the S-S bond to form C2H5SOH and C2H5SO2H. However, no trace of

414

C2H5SOSC2H5 and C2H5SO2SC2H5 was obviously detected in the reaction mixture by

415

GC/MS. It was speculated that the reaction might follow the first pathway.

C2H5SO2-

pH = 9.2

C2H5SO3-

Intensity (μS)

[Fe(VI)] = 0.40 mM

[Fe(VI)] = 0.24 mM

[Fe(VI)] = 0.08 mM [Fe(VI)] = 0 mM

2.0

416

2.5

3.0

3.5

4.0

4.5

5.0

5.5

6.0

Time (min)

417

Fig. 5. The IC chromatograms of products in the oxidization of C2H5SSC2H5 by

418

Fe(VI) at pH = 9.2 in buffer solution. Experimental conditions: [C2H5SSC2H5]0 = 0.4

419

mM, [Fe(VI)] = 0–0.4 mM, reaction time 30 min.

420

The C2H5SO2H is also labile and can further react with Fe(VI) to form

421

C2H5SO3H. The variation in fraction of products at different pH in excess of Fe(VI)

422

could be explained by the oxidizing ability of Fe(VI). The presence of oxygen atom

423

decreases the nucleophilicity of sulfur [20,25]. C2H5SO2H is thus a major oxidized

424

product at pH = 12.0. The appearance of C2H5SO3H in Fig. 4 also demonstrated a

425

much slower rate of the oxidation from C2H5SO2H to C2H5SO3H by Fe(VI) at pH =

426

12.0. As two main final products of C2H5SH degradation, both of C2H5SO2H and

427

C2H5SO3H are not classified either toxic or malodorous chemicals and can therefore 23

428

be discharged into the effluent after neutralization [41,42].

429

In our experiments, a molar ratio of Fe(VI) to C2H5SH for degradation of

430

C2H5SH at pH = 9.2 was determined to be ~1.2, which was slightly less than ~1.4 at

431

pH = 12.0. This result was in consistent with the stoichiometric coefficients predicted

432

by Eqs. (17) and (19). Furthermore, a molar ratio of Fe(VI) to C2H5SH for complete

433

oxidation of C2H5SH to C2H5SO3H was determined to be ~2.2 at pH = 9.2, which was

434

similar to the prediction by Eq. (20). The increase of pH in both reaction mixtures

435

with initial pH = 9.2 and 12.0 without buffer can be described by the two reactions of

436

Eqs. (19) and (20), respectively.

437

7HFeO 4- + 6C2H5SH + 10H2O → 7Fe(OH)3 + 5C2H5SO2H + 1/2C2H5SSC2H5

438

+ 7OH - (17)

439 440 441

2HFeO 4- + C2H5SSC2H5 + 4H2O → 2Fe(OH)3 + 2C2H5SO2H + 2OH (18) 4HFeO 4- + 3C2H5SH + 6H2O → 4Fe(OH)3 + 3C2H5SO2H + 4OH -

442 443

(19) 2HFeO 4- + C2H5SH + 3H2O → 2Fe(OH)3 + C2H5SO3H + 2OH -

444

(20)

24

445 446

Fig. 6. The plausible pathways of C2H5SH oxidation by Fe(VI) in aqueous alkaline

447

solution

25

448

3.4. Comparison with Other Oxidants

449

Table 1. Kinetics, major intermediates and products of C2H5SH with different

450

oxidants in aqueous solution at ambient temperature. pH

Kinetics

O3

4.7

3.0 × 105 M-1 s-1

C2H5SO3H

9.2

2.5 × 106 M-3/2 s-1

C2H5SSC2H5,

12.0

0.9 × 104 M-1 s-1

C2H5SO2H, C2H5SO3H

7.0

~ 107 M-1 s-1 a

6.0

4.0 × 104 M-1 s-1

8.0

4.0 × 106 M-1 s-1

KMnO4

7.0

25.9 M-1 s-1

H2O2

11.0

8.4 M-1 s-1

•OH

6.0

Fe(VI)

chlorine

ClO2

451

a

452

protein with HOCl at pH = 7.0.

453

Major intermediates and

Oxidant

products

C2H5SCl, C2H5SO2Cl C2H5SSC2H5, C2H5SO3H

Reference [41] this study

[37,43]

[44] C2H5SO3H C2H5SOH, C2H5SSC2H5, C2H5SO3H CH3COOH, SO42-

[45] [21,46] [47]

Speculated from the apparent rate constants of sulfur-containing amino acids and

The kinetics, major intermediates and products of C2H5SH degradation by Fe(VI)

454

are compared with other selective and nonselective oxidants (Table 1). As selective

455

oxidants, Fe(VI), O3 , ClO2 and chlorine have high reactivity with C2H5SH. C2H5SH

456

can be rapidly oxidized by Fe(VI) to C2H5SO2H and C2H5SO3H in alkaline solution.

457

The Fe(VI) is finally reduced to a non-toxic chemical of Fe(OH)3 which can act as an

458

effective coagulant in water and wastewater treatment .O3 is an electrophile with high

459

selectivity and unstable in water with a half-life time in the range of seconds to hours

460

depending on the water quality [48,49]. The oxidation of C2H5SH by O3 in acid 26

461

solution is fast resulting in the formation of C2H5SO3H, and the rate constant was

462

found to be independent on pH in the pH range of 0.85 to 4.7 [41]. The rate constant

463

of C2H5SH with ClO2 has been reported to increase exponentially with an increase in

464

pH and estimated as 4.0 × 104 M-1 s-1 and 4.0 × 106 M-1 s-1 at pH = 6 and 8,

465

respectively [44]. However, the dosage of ClO2 is restricted in water treatment

466

because of the blood poison of its reduction products. Reduced sulfur compound (e.g.,

467

sulfhydryl compound and disulfide) could easily be oxidized by chlorine with the rate

468

constants ranged from ~105 M-1 s-1 to ~107 M-1 s-1 in aqueous solution. Chlorine

469

initially oxidizes thiols to yield sulfenyl chloride as an intermediate. And then, the

470

formation of products depends on the chlorination condition, including disulfide,

471

sulfonic acid, and sulfonyl chloride [37]. However, chlorination has potential to form

472

toxic halogenated byproducts.

473

In comparison, the rate constants for KMnO4 and H2O2 are orders of magnitude

474

lower than for Fe(VI), O3 , ClO2 and chlorine. C2H5SH can be oxidized by KMnO4 to

475

C2H5SO3H, but the use of KMnO4 may cause secondary pollution. The rate constants

476

for the oxidation of low molecular weight thiols by H2O2 in aqueous solution were

477

reported in the range of ~10 M-1 s-1 [34]. And the rate constants calculated for thiolate

478

ions were similar and pKa-independent for considered low molecular weight thiols

479

[50]. Oxidation of C2H5SH by H2O2 was proved to be a nucleophilic substitution of

480

C2H5S- on H2O2 with the formation of transient C2H5SOH as the rate determining step

481

[34]. Then, the coupling product of C2H5SSC2H5 would undergo further but much 27

482 483

slower oxidation with H2O2 to give higher oxidation products [21,46]. •OH is a nonselective oxidant with high reactivity and reacts indiscriminately

484

with all kinds of matrix compounds in water [51,52]. It might be an only oxidant that

485

could oxidize C2H5SH to its inorganic form of SO42- in aqueous solution as shown in

486

Table 1. Ma et al. has shown that •OH generated from the modified β-PbO2 anode

487

could react with C2H5SH to produce ethyl radical (•C2H5) and thiyl radical (•HS) by

488

successively attacking on the C-S bond of C2H5SH. The radicals would further react

489

with •OH through a series of reactions and eventually convert to SO42- and

490

CH3COOH, which could be subsequently mineralized [47]. Generally, acidic and

491

neutral pH conditions are required for the generation of •OH in the reactions such as

492

Fenton reaction, photo-catalytic reaction as well as electrochemical system [52,53].

493

However, the acidic and neutral pH is negative for the mass transfer of C2H5SH from

494

gaseous phase to liquid phase in the chemical scrubbing process due to the weak

495

acidity of C2H5SH. With the integrated function of oxidation and coagulation, Fe(VI)

496

is relative more stable in alkaline condition than acidic condition. Therefore, Fe(VI)

497

may be a more suitable oxidant for C2H5SH degradation in the chemical scrubbing

498

technique.

499

4. Conclusions

500

This study demonstrated that C2H5SH could be effectively degraded by Fe(VI) in

501

aqueous alkaline solution. The reaction showed overall 5/2-order behavior at pH =

502

8.0–10.0 and second-order behavior at pH = 10.5–12.0, and the rate constants 28

503

decreased with increased pH. C2H5SH could be finally oxidized to nonvolatile and

504

odorless C2H5SO3H. The stoichiometric ratio of Fe(VI) to C2H5SH for C2H5SH

505

oxidation to C2H5SO3H was determined as 2.0. The plausible mechanism of C2H5SH

506

degradation by Fe(VI) was proposed, in which the two competing reaction pathways

507

for C2H5S• radical were probably responsible for the changes in reaction order at

508

different pH. Compared with other selective and nonselective oxidants, Fe(VI) was

509

revealed to be an efficient and environmental friendly oxidant for C2H5SH

510

degradation. However, challenges still exist in the practical application of Fe(VI) due

511

to the high production cost of solid Fe(VI) products and the instability of Fe(VI)

512

solution. Therefore, the chemical scrubbing process with in situ generation and

513

application of Fe(VI) could be a suitable process to eliminate odors. The odorants

514

could first be absorbed to the liquid phase and then oxidized by the in situ generated

515

Fe(VI) quite rapidly. This may lead to the practical implementation of Fe(VI)

516

technology in odorous gas treatment.

517

Acknowledgements

518

This work was financially supported by State Key Laboratory of Urban Water

519

Resource and Environment (Harbin Institute of Technology) (2014DX08).

520 521

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522

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Graphical abstract

C2H5SH + Fe(VI) pH-dependent kinetics

C2H5SH

Fe(VI)

C2H5SSC2H5 C2H5S •

Fe(VI)

C2H5SO3H

C2H5SOH pH = 10.5 – 12.0

pH = 8.0 – 10.0

 d[Fe(VI)]  k [Fe(VI)][C2H5SH] dt

 d[Fe(VI)]  k [Fe(VI)]3/2[C2H5SH] dt

656 657 658

Highlights

659

C2H5SH can be degraded rapidly by Fe(VI) in aqueous alkaline solution.

660

The reaction kinetics and fraction of intermediates are highly dependent on pH.

661

The reaction shows a 3/2-order dependence on Fe(VI) at pH = 8.0~10.0.

662

Nonvolatile and odorless C2H5SO3H is the final product of C2H5SH degradation.

663

A proposed mechanism successfully explains the pH-dependent reaction order.

664

36