Inhibition and Promotion of Pyrolysis by Hydrogen Sulfide (H2S) and

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Article

Inhibition and Promotion of Pyrolysis by Hydrogen Sulfide (HS) and Sulfanyl Radical (SH) 2

Zhe Zeng, Mohammednoor Altarawneh, Ibukun Oluwoye, Peter Glarborg, and Bogdan Z Dlugogorski J. Phys. Chem. A, Just Accepted Manuscript • DOI: 10.1021/acs.jpca.6b09357 • Publication Date (Web): 25 Oct 2016 Downloaded from http://pubs.acs.org on October 29, 2016

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The Journal of Physical Chemistry

Inhibition and Promotion of Pyrolysis by Hydrogen Sulfide (H2S) and Sulfanyl Radical (SH)

Zhe Zeng,1 Mohammednoor Altarawneh1*, Ibukun Oluwoye,1 Peter Glarborg,2 Bogdan Z. Dlugogorski1

1

School of Engineering and Information Technology, Murdoch University 90 South Street, Murdoch, WA 6150, Australia

2

Department of Chemical Engineering, Technical University of Denmark, DK-2800 Kgs. Lyngby, Denmark

*

Corresponding author:

Phone: +61 89360 7507, Email: [email protected]

1 ACS Paragon Plus Environment


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ABSTRACT

2 3

This study resolves the interaction of sulfanyl radical (SH) with aliphatic (C1-C4)

4

hydrocarbons, using CBS-QB3 based calculations.

5

enthalpies and located the weakest link in each hydrocarbon. Subsequent computations

6

revealed that, H abstraction by SH from the weakest C‒H sites in alkenes and alkynes, except

7

for ethylene, appears noticeably exothermic. Furthermore, abstraction of H from propene, 1-

8

butene and iso-butene displays pronounced spontaneity (i.e., ∆rG° < -20 kJ mol-1 between 300

9

– 1200 K) due to the relatively weak allylic hydrogen bond. On the other hand, an alkyl

10

radical readily abstracts H atom from H2S, with H2S acts as a potent scavenger for alkyl

11

radicals in combustion processes. That is, these reactions proceed in the opposite direction

12

than those involving HS and alkene or alkyne species, exhibiting shallow barriers and strong

13

spontaneity. Our findings demonstrate that, the documented inhibition effect of hydrogen

14

sulfide (H2S) on pyrolysis of alkanes does not apply to alkenes and alkynes.

15

interaction with hydrocarbons, the inhibitive effect of H2S and promoting interaction of SH

16

radical depend on the reversibility of the H abstraction processes. For the three groups of

17

hydrocarbon, Evans-Polanyi plots display linear correlations between the bond dissociation

18

enthalpies of the abstracted hydrogens and the relevant activation energies. In the case of

19

methane, we demonstrated that, the reactivity of SH radicals towards abstracting H atoms

20

exceeds that of HO2 but falls below those of OH and NH2 radicals.

We obtained the C-H dissociation

21


During

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INTRODUCTION

23 24

Hydrogen sulfide (H2S) represents a major impurity in natural gas, and arises in gasification

25

processing of fossil fuels. The pipeline-quality natural gas typically contains about 2000 ppm

26

of sulfur species, in which H2S constitutes the predominant sulfur carrier.1 Although the

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oxidation mechanism of H2S has been studied under atmospheric2 and high pressure3

28

conditions, its interactions with hydrocarbons remain poorly understood. An experimental

29

investigation of Nguyen et al. have indicated that, H2S exhibits an inhibition effect on the

30

pyrolysis of n-octane under pressure of 70 MPa.4,5

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oxidation competition between H2S and CH4, in systems involving the injection of H2S as an

32

additive to CH4/air flames.6,7 Similarly, a recent study by Gerson et al.8 has linked the co-

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oxidation of H2S and CH4 with the H2/O2 chemistry showing that reaction of H2S with the

34

molecular oxygen produces HO2. The latter promotes CH4 oxidation at low and intermediate

35

temperatures.8

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(CH3SH) as a main product in the pyrolysis of alkyl sulfide, which partly origins from the

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direct combination of CH3 and SH radical through negligible activation barrier.9,10

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However, literature lacks detailed understanding of how H2S and hydrocarbons interact with

39

each other, and the H2S inhibition mechanism remains open to speculations.

Likewise, Selim et al. reported the

Additionally, Vandeputte et al. reported the formation of methanethiol

40 41

A pioneering study of Gray et al. investigated the reaction between methyl radical (CH3,

42

produced by photolysis of azomethane) and H2S. The investigators observed that, methyl

43

radical promptly abstracts H atom from H2S, producing a SH radical and a methane molecule

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(see Reaction Ra below).11 Arican et al. further measured the rate constant of Ra as k(T) =

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1.67 × 10-13 exp(-1 054/T) cm3·molecule-1·s-1 within the temperature range of 334 – 432 K,12

46

while an analogous computational study revealed k(T) = 1.13 × 10-16 × T1.2 exp(- 722/T)



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cm3·molecule-1·s-1 for a wider temperature window of 200 – 3000 K.13 On the contrary,

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another experimental work of Perrin et al. discovered that, SH radical can extract one H from

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2-(Z)-C5H10 to form H2S in Reaction Rb, ensuing a reaction rate constant of k(T) = 1.00 × 10-

50

14

exp(-1 159/T) cm3·molecule-1·s-1 for a temperature range of 743 to 772 K.14

51

CH3+ H2S → CH4 + SH

Ra

2-(Z)-C5H10 + ˙SH → CH2CH=CHCH2CH3 + H2S

Rb

52 53

These two sets of reactions proceed in opposite directions, demonstrating that, SH radical

54

may have different effect on alkanes and alkenes. Moreover, to the best of our knowledge,

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literature provides no experimental measurements and theoretical calculations of reactions of

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H2S and/or SH radical with other classes of hydrocarbons.

57 58

While the H abstraction reactions by O/H radical pool act as the initiation step for

59

hydrocarbon oxidation,15 the presence of appreciable concentrations of SH in radical pool at

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elevated temperature could contribute to the overall oxidation mechanism. For this reason,

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this contribution investigates the reactivity of SH radical with a series of hydrocarbons (C1-

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C4), by deploying the density functional theory (DFT) at an adequately high level of theory.

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We cast the reaction sequence according to:

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RH + SH → R···H···SH → R + H2S 65 66

where RH signifies a gas-phase alkane, alkene or alkyne under the C4 chain limit.

67 68

The objectives of this study are to: (i) present the potential energy surfaces for reactions 4 ACS Paragon Plus Environment

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between a SH radical and a series of hydrocarbons, (ii) compute the reaction rate parameters

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in Arrhenius equation over a temperature range from 300 K to 2000 K, (iii) relate the

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dependence of the calculated activation energies with the dissociation enthalpies of the

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relevant C‒H bonds, (iv) compare the reaction activity of a SH radical with other active

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combustion radicals, namely OH, HO2 and NH2 by examining the corresponding bimolecular

74

reactions.

75 76 77

COMPUTATIONAL DETAILS

78 79

Gaussian 09 suite of programs16 facilitated all structural optimisations, and served to

80

calculate enthalpies as well as vibrational frequencies using the complete-basis-set CBS-QB3

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composite method.17 Several studies have already established that, the CBS-QB3 method

82

offers accurate calculations of geometries and activation enthalpies of hydrogen abstraction

83

reactions.18-20 For instance, a concise study of Pokon et al. applied this method to calculate a

84

series of hydrocarbon deprotonation reactions in gas phase, with the results indicating a

85

maximum deviation of 6.2 kJ·mol-1, relative to the corresponding experimental values.21 We

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further validated the accuracy of CBS-QB3 method and corroborate this result in the

87

discussion part of the present article.

88 89

In the calculations, the absence of imaginary frequencies in all reactant and product species

90

indicated true energy minima. In contrast, the computed transition structures retained one,

91

and the only one imaginary frequency along the specific reaction coordinate. Since the effect

92

of hindered rotors in reactants and transition structures cancels each other, we treated all

93

hindered rotors as harmonic oscillators.

In addition, for selected reactions, we applied



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intrinsic reaction coordinate (IRC) calculation to confirm the reaction pathways.

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The ChemRate program22, based on the classical transition state theory (TST)23, afforded

97

estimation of the reaction rate constants. As implemented in the KiSTheIP code, a one-

98

dimensional asymmetrical Eckart barrier accounted for quantum tunneling effects on

99

computed rate coefficients.24,25 The reaction rate parameters were fitted to the Arrhenius

100

equation in the form of k(T) = A·Tn·exp(-Ea/RT), over a temperature range of 300 K to 2000

101

K. Finally, Chemkin-Pro26 program served to obtain the equilibrium species concentrations

102

for selected systems.

103 104 105

RESULTS AND DISCUSSION

106 107

Overview of Labile H Abstraction Sites in C1-C4 Hydrocarbons

108 109

In general, the titled reactions involve the following process:

110

RH + SH → R···H···SH → R + H2S 111 112

in which SH radical abstracts single H from an hydrocarbon, and as a result, produces a

113

radical R and H2S via the transition state [R···H···SH] structure. Most hydrocarbons display

114

different sites for hydrogen abstraction by SH. For example, the SH radical can extract H

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from primary or secondary C‒H site in propane, resulting in a different radical R and a

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transition structure. The main interest lies in elucidating the most favorable routes during

117

combustion processes. 6 ACS Paragon Plus Environment

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118 119

For this reason, we calculated the bond dissociation enthalpies (BDH) to locate the weakest

120

C‒H bonds in all species studied in this paper. In subsequent computations, we assumed that,

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SH radical preferentially attacks the weakest C‒H bond in the hydrocarbons. Further

122

calculation validates that H abstraction from weakest C-H site in hydrocarbon is featured

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with lower activation barrier in case of propane. Our computed BDH values have been

124

compared with analogous experimental and the theoretical estimations summarised by Luo.27

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Table 1 demonstrates agreements between the recommended values of Luo27 and the present

126

results, within the expected accuracy of the CBS-QB3 method.

127 128

Table 1.

129 130

In the case of propane and n-butane, the secondary carbon displays the weakest C-H bond,

131

while for iso-butane, the weakest C-H bond is that on associated with the tertiary carbon. For

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the alkenes and alkynes, the most vulnerable C‒H bond exists on the saturated carbon sites.

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In brief, the following sequence reflects the strength of the BDH of C‒H bonds:

134 135

C‒H on unsaturated carbon > C‒H on saturated carbon (primary > secondary > tertiary)

136 137

Table 2 presents all hydrogen abstraction reaction considered in the present study. These

138

reaction involve the weakest C-H bond of the target molecule.

139 140

Table 2.

141 142



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H Abstraction from Alkanes

144 145

In this section, we discuss the reactions of SH radical with alkanes, namely, methane, ethane,

146

propane, n-butane and iso-butane, designated by R1 through R6 in Table 2. Figure 1 portrays

147

all the relevant transition structures. For example, in the structure of transition state of R1,

148

the distance between the C atom and the dissociated H atom increases to 1.64 Å (0.164 nm),

149

whereas the separation between the abstracted H and S atoms corresponds to 1.44 Å.

150

Examination of Fig. 1 reveals similar distances for other transition structures.

151

Figure 1.

152 153 154

Figure 2 depicts the potential energy surface for reactions between all C1-C4 alkanes and the

155

SH radical. For methane (R1), the activation enthalpy corresponds to 65.9 kJ·mol-1. While in

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R2, reaction of SH radical with ethane molecule requires a relatively smaller barrier height of

157

45.7 kJ·mol-1. This concurs with the trend of comparative BDH of methane (440.9 kJ·mol-1)

158

and ethane (420.5 kJ·mol-1). In the case of propane, H abstraction from primary carbon site

159

(BDH of 428.5 kJ·mol-1) in R3 requires a higher activation barrier of 46.6 kJ·mol-1 when

160

compared to that of the weaker secondary links (BDH of 413.9 kJ·mol-1) in R4 (29.1 kJ·mol-

161

1

162

which is consistent with our previous assumption that SH radicals prefer to attack the weakest

163

C‒H bond in the hydrocarbons with lower activation barriers. In the remainder of the paper,

164

we focus on the most favourable channel for H abstraction from each molecule.

). Clearly, H abstraction from a weaker C‒H bond results in a lower activation enthalpy,

165 166

Figure 2.

167


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As demonstrated in Fig. 2 for alkanes, the enthalpy of the separated reaction products (R +

169

H2S) exceeds that of separated reactants (RH + SH). The endothermic condition implies that,

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the reverse reactions, i.e., R1b, R2b, and R3b, prevail over forward Reactions R1, R2, and

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R3. This agrees well with the experimental and calculation results available in literature11-13,

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in which H2S reacts with methyl radical to form methane as the primary product. Table 3

173

compares the calculated kinetic parameters of R1b with the analogous experimental

174

measurements. The reaction rate corresponds to 1.66 × 10-14 cm3·molecule-1·s-1 at 423 K, in

175

good agreement with the equivalent experimental value11 of 1.45 × 10-14 cm3·molecule-1·s-1.

176

The agreement validates the present computational approach to estimate the kinetic

177

parameters of similar reactions that have no experimentally-measured values of rate

178

constants.

179 180

Table 3.

181 182

For propane, n-butane and iso-butane (Reaction R4, R5 and R6), the relative enthalpies of the

183

TS structures lie between that of the reactants and products. In effect, this means that, the

184

reverse reactions proceed without an enthalpy barrier. In Fig. 3, we present the results of the

185

IRC calculation to confirm the transition state structures and the reaction pathways. All

186

displayed points are calculated at a fixed C‒H‒S distance around the obtained TS, featuring

187

one imaginary frequency. Although, the structures rest within the saddle point of each curve,

188

the consecutive enthalpies increase as they approach the product. This behaviour necessitates

189

the variational transition state theory (VTST) calculations to obtain the reaction rate

190

constants. For this purpose, we employed five points adjoining the product and minimised

191

the reaction rate constant as a function of temperature and reaction coordinate. Due to the

192

absence of a transition structure, we are unable to evaluate the rate parameters for reverse



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193

reactions directly. Herein, we calculated an equilibrium constant with Chemkin26 at each

194

temperature to derive backward reaction rate using of the rate of the forward reaction we

195

obtained with VTST (see supporting information for detail).

196 197

Figure 3.

198 199

As these reactions are considerably endothermic with a shallow or no reverse barrier, the

200

backward reaction dominates, with insignificant formation of H2S and CH3. That is, alkyl

201

radicals (CH3, C2H5, C3H7 and C4H9) readily abstract H atoms from H2S to produce alkane

202

molecules and SH radicals. For C2H5 and prim-C3H7, the trivial activation barriers amount to

203

4.9 kJ·mol-1 and 2.9 kJ·mol-1, respectively at 298.15 K. For the remaining alkyl radicals, the

204

reactions proceed via barrier-less processes.

205 206

To gain further insights into the equilibrium condition of R1, we performed simple

207

equilibrium calculations with Chemkin-Pro26. The procedure required the thermochemical

208

data of the reactants (relatively to the forward path, i.e. CH4 and SH) and the products (CH3

209

and H2S) as derived from their respective enthalpy of formation (NIST webbook), vibrational

210

frequencies and rotational constants (computed by CBS-QB3) with ChemRate22. In addition,

211

we also compared the results with using thermodynamic data from Burcat’s database.28

212

Figure 4 displays the equilibrium quantities of the reacting species for an initial concentration

213

of 25 % v/v CH4, and 25 % v/v SH in nitrogen balance, at different temperatures (300 – 2000

214

K), calculated with our derived thermodynamic data and that of Burcat (see supporting

215

information), respectively. The difference between two sets of data is within 6 %， which

216

also validates the accuracy of the derived thermochemical data used in this work. The

217

equilibrium concentration of reactants significantly exceeds those of the products, implying


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that, the backward reaction plays a more significant role during H abstraction from alkanes.

219

To this end, one can easily observe that the presence of H2S hinders the pyrolysis of CH4,

220

concurring with the inhibition effect on n-octane.4,5 However, under oxidative conditions,

221

the production of HO2 (i.e., a major oxidation carrier in the low – intermediate temperature

222

windows) during the early stages of H2S oxidation (H2S + O2 → HS + HO2) may overshadow

223

the overall inhibition activity of H2S during CH4/H2S co-oxidation scheme. This has been

224

manifested by kinetic sensitivity analysis8 and a noticeable reduction in autoignition delay

225

time upon the addition of 1% H2S to methane at pressures ranging from 30-80 bar and

226

temperatures from 930-1050 K.

227

Figure. 4

228 229

Table 4 reports the Arrhenius parameters for the forward and the reverse reactions of H

230

abstractions from the weakest C‒H bond in C1-C4 alkanes. The rate parameters for Reactions

231

R4, R5 and R6 come from the VTST calculations, while those of R4b, R5b and R6b are

232

derived with equilibrium constant Kc and forward reaction rate kf at each temperature,

233

respectively. The classical TST formalism yielded the remaining rate parameters.

234 235

Table 4.

236 237 238

H Abstraction from Alkenes and Alkynes

239 240

In Fig. 5, we present the calculated potential energy surface for reactions involving ethylene,

241

propene, 1-butene, iso-butene, allene, propyne and 1-butyne.

242

comprise the abstraction of H atoms from the weakest C‒H bond in each hydrocarbon.


All illustrated reactions


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243

Hydrogen abstraction from ethylene by SH differs significantly from other systems, with the

244

reaction endothermicity of 78.7 kJ·mol-1, and the transition state lying 3.4 kJ·mol-1 below the

245

separated products.

246

reaction pathway and it demonstrated a similar pattern with propane, n-butane and iso-butane,

247

as illustrated in Fig. 3. For the remaining species, the reactions are all exothermic and

248

proceed with plausible activation barriers. As opposed to alkanes, SH radical can extract one

249

H from the weak C‒H sites on alkenes/alkynes (except for ethylene) to produce H2S. This is

250

consistent with the experimental work of Perrin et al.14, in which the authors observed the

251

formation of H2S as a result of H atom migration from 2-(Z)-C5H10 towards SH radical.

Consequently, an IRC calculation was conducted to elucidate the

252

Figure 5.

253 254 255

The H abstraction process in 1-butene incurs a trivial activation enthalpy of 1.8 kJ·mol-1 with

256

reaction exothermicity of 39.4 kJ·mol-1. Noticeably, the barrier heights rise to 16.7 kJ·mol-1

257

and 13.0 kJ·mol-1, for propene and 1-butyne, respectively. But for iso-butene, allene and

258

propyne, the reactions involve activation enthalpies of 39.8 kJ·mol-1, 35.5 kJ·mol-1 and 34.7

259

kJ·mol-1, in that order. All these values reside within the accuracy limit of the computational

260

methodology (CBS-QB3). Table 5 assembles the fitted Arrhenius parameters for Reaction R7

261

through R13. We obtained the rate parameters for Reaction R7 from VTST, and used TST for

262

the other reactions.

263 264

Table 5.

265 266

Validation of CBS-QB3 Calculations

267


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Although, the experimental data for H abstraction from hydrocarbons by SH radical are

269

scarce, we conducted three sets of comparison to confirm the accuracy of the calculated CBS-

270

QB3 results. Firstly, as shown in Table 1, we have computed the BDH values for different H-

271

abstraction site in aliphatic hydrocarbons, and compared them with analogous experimental

272

and theoretical estimations summarised by Luo27. The calculated results match the

273

recommended BDH values, but it appears that, the molecular complexity of higher

274

hydrocarbons leads to larger errors in BDHs. We computed the absolute mean error to be 5.9

275

kJ·mol-1, in good agreement with the results of another work of Pokon et al.21, who adopted

276

the same basis set to calculate a series of gas-phase deprotonation reactions with mean

277

absolute deviation of 6.2 kJ·mol-1 from experimental values.

278 279

Secondly, we compared the reaction rate of Ra (CH3 + H2S → CH4 + SH) with the

280

experimental data. As mentioned earlier, this reaction had previously been studied both

281

experimentally and theoretically11-13. Figure 6 contrasts our calculated rate constants with the

282

literature values. Our results compare well with the experimental data that were obtained for

283

a relative narrow temperature range (334 – 432 K). On the other hand, the Arrhenius

284

expression of Mousavipour et al.13, computed at a lower level of theory (MP2/6-311+G(d,p)),

285

overestimates the rate coefficient at lower temperatures.

286

constants of R1 derived from the thermochemical data calculated in this work deviate by only

287

6% of analogous values calculated based on Burcat’s database.28 This serves as a validation

288

for the accuracy of calculated vibrational frequencies and rotational constant within the

289

adapted CBS-QB3 methodology.

290

temperature dependency of Gibbs free energy (∆rG°) for all reactions (from 300 K – 1200 K).

291

The decreasing trend of Gibbs free energy indicates that all considered H abstraction

292

reactions become more spontaneous as the temperature increases.

Additionally, the equilibrium

Figure S1 in supporting information demonstrates the



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293

Figure 6.

294 295 296

Thirdly, we examined Reaction Rb (2-(Z)-C5H10 + SH → CH2CH=CHCH2CH3 + H2S, see

297

Fig. S2 in supporting information) to confirm the accuracy of our calculation when compared

298

with the experimental data of Perrin et al.14

299

Arrhenius equation within the temperature range of 300 – 2000 K to obtain k(T) = 6.51 × 10-

300

14

301

results fitted between 743 – 772 K to k(T) = 1.0 × 10-14 exp(-1 159/T) cm3·molecule-1·s-1, with

302

a relatively larger pre-exponential factor.

303

aforementioned reaction, and the supporting information provides the relevant potential

304

energy surface.

305

Bearing in mind plausible source of errors in the experimental work as well the accuracy

306

benchmark of the adopted theoretical methodology, the difference depicted in Fig. 7 remains

307

within an order of magnitude. A plausible source for the discrepancy in Figure 7 might stem

308

from treating all hindered rotors in the (2-(Z)-C5H10) as harmonic oscillators. It is a very

309

daunting task to accurately account for all coupled internal rotations in the rather complex

310

molecular structure of (2-(Z)-C5H10).

Our calculated results were fitted to the

exp(-1 219/T) cm3·molecule-1·s-1. This equation somewhat overestimates the experimental

Figure 7 assesses the Arrhenius plot of the

Our calculated data seem to somewhat overestimate the reaction rate.

311 312

Figure 7.

313 314

Relationship between BDH of Abstracted C‒H Bond and Activation

315

Enthalpy of SH + Hydrocarbon Reaction

316 317

To explain the different behaviour of H abstraction from alkane, alkene and alkyne by SH 14 ACS Paragon Plus Environment

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radical, we introduced the Evans-Polanyi plots as illustrated in Fig. 8. The abstraction rate

319

(represented by activation enthalpy) generally depends on the strengths of the dissociated C‒

320

H bond (as reflected by the corresponding BDH) and the activity of the attacking radical

321

(SH). The figure compares the linear correlations between activation enthalpy ∆E and BDH

322

for each group of hydrocarbons.

323

Figure 8.

324 325 326

The orbital hybridisation in double and triple-bonded carbon leads to strong C‒H bonds on

327

the unsaturated carbon atoms. For example, all H atoms of ethylene connect to unsaturated

328

carbon atoms (=CH2), which results in significantly higher BDH (462.8 kJ·mol-1) if compared

329

with allylic C-H bonds in propene (359.3 kJ·mol-1). This explains the outlier position of

330

ethylene among the studied alkene molecules (see Fig. 8), and elucidates the reason for the

331

C2H3 radical abstract H from H2S, via a barrierless pathway.

332 333

As illustrated in Fig. 8, the weakest C-H bond in alkanes is much stronger than that in

334

alkenes/alkynes. Generally, the extra orbital overlap on unsaturated C weakens the C‒H

335

bond on the vicinal saturated carbons, resulting in lower BDH for alkenes and alkynes (i.e.,

336

on saturated C site) compared to alkanes. Based on our kinetic analysis, we report that, the

337

forward reaction predominates for alkenes and alkynes, whereas the backward process

338

governs the reactions involving alkanes.

339 340

Furthermore, since the reaction rate relies on the strength of the dissociated C‒H bond and

341

the reaction activity of the SH radical, we calculated the BDH of S-H in H2S molecule as

342

380.4 kJ·mol-1. This value forms a distinct boundary between the set of reactions that



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

343

proceed in a forward direction and those that advance in the reverse direction. On the left

344

side (Fig. 8), BDHs of alkenes and alkynes are smaller than that of S-H. This means that, SH

345

can abstract H from these species. On the contrary, the C‒H bond in alkanes remain much

346

stronger than the S‒H bond in H2S, resulting in species to the right of 380.4 kJ·mol-1

347

abstracting H from H2S, to form alkane molecules.

348 349 350

Activity of SH Radical Compared with Those of OH, NH2 and HO2

351 352

In order to gain further insights into the role of SH radical in combustion processes, we

353

explored the reactivity of SH radical, by comparing it with that of NH2, OH and HO2 from

354

literature. Figure 9 contrasts the rate constants for H abstraction from methane by the four

355

radicals.

356

overlapping rate constants. Theoretical calculations of Mebel and Lin yielded 60.2 kJ·mol-1

357

as the activation energy for the reaction CH4 + NH2 → CH3 + NH3,29 i.e., very close to the

358

value of 61.2 kJ·mol-1 calculated herein for analogous reaction involving SH radical.

359

Abstraction by OH30 and HO231 radicals incurs the fastest and the slowest reaction rate,

360

respectively. The reaction of CH4 + NH2 → NH3 + CH3 proceeds predominantly in the

361

forward direction, in contrast to similar abstraction by the SH radical. Comparably, HO2

362

radical seems relatively inactive with a reported activation barrier of 87.9 kJ·mol-1. Thus, we

363

conclude that, the radical activities follow the order of OH > NH2 > HS > HO2.

Interestingly, H abstraction from methane by SH and NH2 radicals display

364 365

Figure 9.

366 367

This ordering follows the strengths of the freshly formed O‒H, N‒H and S‒H bonds. For 16 ACS Paragon Plus Environment

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368

example, the S‒H bond is weaker than O‒H and N‒H bonds and hence OH and NH2 remain

369

more effective in abstracting H atom from the hydrocarbon chain. Along the same line of

370

enquiry, the O‒H bond in H2O2 constitutes a weaker bond justifying its lower reactivity.

371 372 373

CONCLUSIONS

374 375

This contribution has reported the thermokinetic parameters for reactions of SH radicals with

376

a series of C1-C4 hydrocarbons. The inhibition effect of H2S on pyrolysis of alkanes stems

377

from the facile abstraction of H atom from H2S by alkyl radicals, exposing H2S as an

378

effective scavenger of alkyl radicals. The presence of weaker allylic C‒H bonds in alkenes

379

and alkynes forces the overall reaction to proceed in the forward direction, i.e., in the

380

direction of forming the H2S and alkenyl and alkynyl radicals. A linear relationship exists

381

between activation barriers and bond dissociation enthalpies of the attacked C‒H sites. BDH

382

of H2S of 380.4 kJ·mol-1 separates the species whose radicals abstract H from H2S, from

383

those whose hydrogen atoms are abstracted by HS.

384

compared to that of OH, HO2 and NH2 radicals follows the order of OH > NH2 > HS > HO2.

The reactivity of the HS radical

385 386 387

ASSOCIATED CONTENT

388 389

Supporting Information

390 391

The Supporting Information is available free of charge on the ACS Publications website at

392

DOI:



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

393 394

Coordinates and calculated enthalpy values for all structures, thermodynamic data

395

calculated for R1, Gibbs free energy for all reactions at high temperature calculated in

396

this work

397 398 399

AUTHOR INFORMATION

400 401

Corresponding Author

402 403

* Email: [email protected], Phone: +61 8 9360 7507

404 405

Notes

406 407

The authors declare no competing financial interest.

408 409 410

ACKNOWLEDGEMENT

411 412

This study has been supported by grants of computing time from the National Computational

413

Infrastructure (NCI) Australia and from the Pawsey Computing Centre in Perth as well as

414

funds from the Australian Research Council (ARC).

415

University for postgraduate research scholarships.

Z. Z. and I. O. thank Murdoch

416 417


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418


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2. Zhou, C.; Sendt, K.; Haynes, B. S. Experimental and kinetic modelling study of H2S

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oxidation. Proc. Combust. Inst 2013, 34, 625-632.

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3. Song, Y.; Hashemi, H.; Christensen, J. M.; Zou, C.; Haynes, B.; Marshall, P.; Glarborg, P.

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5. Nguyen, V. P.; Burklé-Vitzthum, V.; Marquaire, P. M.; Michels, R. Pyrolysis mechanism of

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decomposition of dimethyl disulfide. J. Phys. Chem. A, 2010, 114, 10531-10549.

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11. Gray, P.; Herod, A. A.; Leyshon, L. J. Abstraction of hydrogen from hydrogen sulfide by

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methyl radicals. Can. J. Chem. 1969, 47 , 689-690.

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12. Arican, H.; Arthur, N. Reactions of methyl radicals. V. Hydrogen abstraction from

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hydrogen sulfide. Aust. J. Chem. 1983, 36, 2195-2202.

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13. Mousavipour, S. H.; Namdar-Ghanbari, M. A.; Sadeghian, L. A theoretical study on the

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kinetics of hydrogen abstraction reactions of methyl or hydroxyl radicals with hydrogen

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sulfide. J. Phys. Chem. A 2003, 107, 3752-3758.

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14. Perrin, D.; Richard, C.; Martin, R. H2S-promoted thermal isomerization of cis-2-pentene

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to 1-pentene and trans-2-pentene around 800 K. Int. J. Chem. Kinet. 1988, 20, 621-632.

454

15. El Marrouni, K.; Abou-Rachid, H.; Kaliaguine, S. Density functional theory kinetic

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16. Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J.

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463

18. Casasnovas, R.; Frau, J.; Ortega-Castro, J.; Salvà, A.; Donoso, J.; Muñoz, F.

464

Simplification of the CBS-QB3 method for predicting gas-phase deprotonation free energies.

465

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19. Vandeputte, A. G.; Sabbe, M. K.; Reyniers, M. F.; van Speybroeck, V.; Waroquier, M.;

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20. Vandeputte, A. G.; Reyniers, M. F. A theoretical study of the thermodynamics and kinetics

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of small organosulfur compounds. Theor. Chem. Acc. 2009, 123, 391-412.

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21. Emma K. Poken; Matthew D. L.; Steven F.; George C. S. Comparison of CBS-QB3,

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CBS-APNO, and G3 predictions of gas phase deprotonation data. J. Phys. Chem. A 2001,

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105, 10483-10487.

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22. Mokrushin V.; Bednov V.; Tsang, W.; Zachariah M.; Knyazev V. ChemRate, NIST:

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23. Truhlar, D. G.; Garrett, B. C.; Klippenstein, S. J. Current status of transition-state theory.

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J. Phys. Chem. A 1996, 100, 12771-12800.

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24. Eckart C. The penetration of a potential barrier by electrons. Phys. Rev. 1930, 35 (11),

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25. Canneaux, S.; Bohr, F.; Henon, E. KiSThelP: a program to predict thermodynamic

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26. Reaction Design, San Diego, CHEMKIN-PRO, 15131; 2013.

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27. Luo, Y.-R. Handbook of bond dissociation energies in organic compounds. Taylor and

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28. Burcat, A. Ideal gas thermodynamic data in polynomial form for combustion and air

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pollution use (http://garfield.chem.elte.hu/Burcat/burcat.html).

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29. Mebel, A. M.; Lin, M. C. Prediction of absolute rate constants for the reactions of NH2

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with alkanes from ab initio G2M/TST calculations. J. Phys. Chem. A 1999, 103, 2088-2096.

490

30. Atkinson, R. Kinetics of the gas-phase reactions of OH radicals with alkanes and

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cycloalkanes. Atmos. Chem. Phys. 2003, 3, 2233-2307.

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31. Aguilera-Iparraguirre, J.; Curran, H. J.; Klopper, W.; Simmie, J. M. Accurate benchmark



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calculation of the reaction barrier height for hydrogen abstraction by the hydroperoxyl radical

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from methane: implications for CnH2n+2 where n = 2 → 4. J. Phys. Chem. A 2008, 112, 7047-

495

7054.


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Table 1. Comparison of the calculated BDH with the literature values for all hydrocarbons studied in this work.27 Hydrocarbon

The broken bonds

BDH calculated

Recommended BDH

Absolute

species

(in bold)

in this work

[27]

error

(kJ·mol-1)

(kJ·mol-1)

(kJ·mol-1)

Methane

CH4

440.9

439.3±0.4

1.6

Ethane

CH3CH3

425.5

420.5±1.3

5.0

CH3CH2CH3

428.5

422.2±2.1

6.3

CH3CH2CH3

413.9

410.5±2.9

3.4

CH3CH2CH2CH3

426.4

421.3

5.1

CH3CH2CH2CH3

414.8

411.1±2.2

3.7

(CH3)2CHCH3

428.1

419.2±4.2

8.9

(CH3)3CH

406.6

400.4±2.9

6.2

CH2=CH2

462.8

465.3±3.3

2.5

CH3CH=CH2

466.7

464.8

1.9

CH3CH=CH2

446.6

N/A

N/A

CH3CH=CH2

359.3

368.6±2.9

9.3

CH2=CHCH2CH3

418.9

410.5

8.4

CH2=CHCH2CH3

346.1

350.6

4.5

CH2=CHCH2CH3

447.6

N/A

N/A

CH2=CHCH2CH3

465.9

N/A

N/A

CH2C(CH3)CH3

371.3

362.8±2.5

8.5

CH2C(CH3)2

469.6

N/A

N/A

CH2=C=CH2

379.3

371.1±12.6

8.2

Propane

n-Butane

iso-Butane

Ethylene

Propene

1-Butene

iso-Butene

Allene



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 24 of 39

Propyne

CH≡CCH3

382.6

372.0±4.2

10.6

1-Butyne

CH≡CCH2CH3

362.3

355.6

6.7

Mean Absolute Error

5.9


Page 25 of 39

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Table 2. Reactions studied in this work, C in bold indicates the H abstraction site(s) for each hydrocarbon. R1b, R2b and R3b denote the reverse reactions of R1, R2 and R3. Reactant

Reaction CH4 + SH → CH3 + H2S

R1

CH3 + H2S →CH4 + SH

R1b

C2H6 + SH → CH2CH3 + H2S

R2

CH2CH3 + H2S → C2H6 + SH

R2b

C3H8 + SH → CH2CH2CH3 + H2S

R3

CH2CH2CH3 + H2S →C3H8 + SH

R3b

C3H8 + SH → CH(CH3)2 + H2S

R4

CH(CH3)2 + H2S → C3H8 + SH

R4b

C4H10 + SH → CH3CHCH2CH3 + H2S

R5

CH3CHCH2CH3 + H2S → C4H10 + SH

R5b

C4H10 + SH → C(CH3)3 + H2S

R6

C(CH3)3 + H2S → C4H10 + SH

R6b

Ethylene

C2H4 + SH → CH=CH2 + H2S

R7

Propene

C3H6 + SH → CH2=CHCH2 + H2S

R8

1-Butene

C4H8 + SH → CH2=CHCHCH3 + H2S

R9

iso-Butene

C4H8 + SH → CH2=C(CH2CH3) + H2S

R10

Allene

C3H4 + SH → CH2=C=CH + H2S

R11

Propyne

C3H4 + SH → C≡CCH2 + H2S

R12

1-butyne

C4H6 + SH → CH≡CCHCH3

R13

Methane

Ethane

Propane

n-Butane

iso-Butane



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 26 of 39

Table 3. Arrhenius parameters fitted between 334 K and 432 K for Reaction R1b; k(T) = A·exp(-Ea/RT) for both experimental measurements12 and theoretical calculations. A

Ea

(cm3·molecule-1·s-1)

(kJ·mol-1)

Experiment

1.67 × 10-13

8.8

[12]

TST with Eckart

6.76 × 10-13

13.3

This work

Source

Method


Page 27 of 39

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Table 4. Arrhenius parameters fitted between 300 and 2000 K for alkane + SH radical; k(T) = A·Tn· exp(-Ea/RT). A

Ea n

Reaction (cm3·molecule-1·s-1)

(kJ·mol-1)

R1: CH4 + SH → CH3 + H2S

7.78 × 10-22

3.02

66.3

R2: C2H6 + SH → CH2CH3 + H2S

4.37 × 10-22

3.41

42.2

R3: C3H8 + SH → CH2CH2CH3 + H2S

8.51 × 10-22

3.39

43.2

R4: C3H8 + SH → CH(CH3)2 + H2S

5.25 × 10-18

1.79

34.6

R5: C4H10 + SH → CH3CHCH2CH3 + H2S

3.26 × 10-20

2.53

31.3

R6: C4H10 + SH → C(CH3)3 + H2S

1.56 × 10-17

1.94

24.3

R1b: CH3 + H2S →CH4 + SH

2.15 × 10-22

3.15

3.4

R2b: CH2CH3 + H2S → C2H6 + SH

5.89 × 10-23

3.06

1.1

R3b: CH2CH2CH3 + H2S →C3H8 + SH

5.25 × 10-22

2.74

0.4

R4b: CH(CH3)2 + H2S → C3H8 + SH

7.64 × 10-21

3.61

8.7

R5b: CH3CHCH2CH3 + H2S → C4H10 + SH

6.86 × 10-13

0.03

12.1

R6b: C(CH3)3 + H2S → C4H10 + SH

1.03 × 10-12

0.03

9.6



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 28 of 39

Table 5. Arrhenius parameters fitted between 300 and 2000 K for alkene + SH radical and alkyne + SH radical; k(T) =A·Tn·exp(-Ea/RT). A

Ea n

Reaction (cm3·molecule-1·s-1)

(kJ·mol-1)

R7: C2H4 + SH → CH=CH2 + H2S

2.96 × 10-25

3.31

81. 3

R8: C3H6 + SH → CH2=CHCH2 + H2S

2.00 × 10-24

3.79

9.9

R9: C4H8 + SH → CH2=CHCHCH3 + H2S

2.19 × 10-23

3.40

0.4

R10: C4H8 + SH → CH2=C(CH2CH3) + H2S

2.69 × 10-22

3.32

36.5

R11: C3H4 + SH → CH2=C=CH + H2S

2.51 × 10-22

3.37

30.2

R12: C3H4 + SH → C≡CCH2 + H2S

2.24 × 10-21

3.36

29.1

R13: C4H6 + SH → CH≡CCHCH3

1.10 × 10-22

3.32

8.01


Page 29 of 39

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Figure 1. Transition state structures for H abstraction from the weakest C‒H bond of the studied hydrocarbons. Distances are in Å.



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Figure 2. Potential energy surface of C1-C4 alkanes + SH radical, computed at CBS-QB3 level of theory. H atom is abstracted from weakest C-H, as illustrated by the product radicals. Values of relative enthalpy and Gibbs free energy (in italic and brackets) at 298.15 K, are in kJ·mol-1.


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Figure 3. IRC calculation for reaction pathway of SH + propane/n-butane/iso-butane. Red dots indicate the transition structure identified at saddle point. Green dots denote the structures adopted in VTST calculations.



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Figure 4. The equilibrium amounts of reactants (CH4, SH) and products (CH3, H2S) for an initial mixture of 25 % v/v CH4, 25 % v/v of SH and 50 % v/v of dilution N2 at different temperatures. Calculations are conducted at constant pressure (1 atm) and temperature for each point.


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Figure 5. Potential energy surface for selected alkene/alkyne + SH radical, computed at CBSQB3 level of theory. H atom is abstracted from weakest C-H, as illustrated by the product radicals. Values of relative enthalpy and Gibbs free energy (in italic and brackets) at 298.15 K

are

in


kJ·mol-1.


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Figure 6. Comparison of rate coefficients obtained in this work (CBS-QB3 basis set) with experimental and theoretical (MP2/6-311++G(d,p)) values for Reaction Ra: CH3+ H2S → CH4 + SH.


Page 34 of 39

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Figure 7. Comparison of rate coefficients obtained in this work (CBS-QB3 basis set) with experimental values for Reaction Rb: 2-(Z)-C5H10 + SH → CH2CH=CHCH2CH3 + H2S.



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Figure 8. Evans-Polanyi plots (activation enthalpy of SH + hydrocarbons versus BDHs of the weakest C-H bonds in hydrocarbon molecules). The black border denotes the BDH of H2S at 380.4 kJ·mol-1. All numbers are in kJ·mol-1 at 298.15 K.


Page 36 of 39

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Figure 9. Rate constants for H abstraction from CH4 by SH, NH2, HO2 and OH radicals, units of k are in cm3·molecule-1·s-1. Results are taken from this work (SH + CH4) and the literature (NH2/HO2/OH + CH4).



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TOC/ABSTRACT ART


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Graphical Abstract 99x81mm (96 x 96 DPI)

ACS Paragon Plus Environment

Inhibition and Promotion of Pyrolysis by Hydrogen Sulfide (H2S) and

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