Oxidation chemistry of carbon disulfide (CS and its ... · Kamal Siddique, Arif Abdullah, Jomana Al-Nu’airat, Sidra Jabeen, Dr Anam Saeed, Dr Juita and Dr Jakub Skut. Courtesy of

Oxidation chemistry of carbon disulfide (CS2)

and its interaction with hydrocarbons in

combustion processes

Zhe Zeng, BEng

This thesis is presented for the degree of

Doctor of Philosophy

School of Engineering and Information Technology,

Murdoch University, Western Australia

2017

i

Statement of originality

I declare that this thesis is my own account of my research and contains as its main content

work which has not previously been submitted for a degree at any tertiary education institution.

The thesis contains no material previously published or written by another person, except

where due reference has been made in text.

Zhe Zeng

June, 2017

ii

Supervisory statement

We, the undersigned, attest that Higher Research Degree candidate, Zhe Zeng, has devised and

synthesised the experimental program, conducted experiments, analysed data, performed

computational quantum-mechanical calculations and has written all papers included in this

thesis. Professor Bogdan Z. Dlugogorski and Dr Mohammednoor Altarawneh provided the

necessary advice on the experimental program, project direction and assisted with the editing

of the papers, consistent with normal supervisors-candidate relations.

Professor Bogdan Z. Dlugogorski

June, 2017

Dr Mohammednoor Altarawneh

June, 2017

iii

Acknowledgment

I would like to thank my supervisors Professor Bogdan Z. Dlugogorski and Dr Mohammednoor

Altarawneh for their unwavering and exceptional support. Their understanding, mentorship,

encouragement, dedication and magnanimity have provided a good fundamental driving force

for the present study. Thank you again, for your quick correspondence, corrections and

gestures of support.

I am grateful to Murdoch University for the award of a postgraduate research scholarship which

provided a valuable financial support for this research. This study has also been funded by the

Australian Research Council (ARC), with grants of computing time from the National

Computational Infrastructure (NCI) and the Pawsey Supercomputing Centre in Perth,

Australia.

I acknowledge my gratitude to my fellow student colleagues and staffs of the Fire Safety and

Combustion Kinetics Research Laboratory: Ibukun Oluwoye, Niveen Assaf, Nassim Zeinali,

Kamal Siddique, Arif Abdullah, Jomana Al-Nu’airat, Sidra Jabeen, Dr Anam Saeed, Dr Juita

and Dr Jakub Skut. Courtesy of you all, I experienced an exciting and fun-filled years of

computation and experiments.

Special thanks to my wife Anqi Wang, who supported me through the most difficult period of

my research. This experience is the most memorable and valuable days in my life.

iv

Abstract

This thesis presents a series of scientific studies exploring the oxidation chemistry of carbon

disulfide (CS2) and its interaction with hydrocarbons in combustion systems. The results

illustrate the extreme flammability of CS2 even at low temperature (self-ignition temperature

at 363 K under ambient pressure). The thesis proposes a comprehensive oxidation mechanism

that works over a wide range of atmospheric and combustion conditions with and without

moisture. The thesis also provides experimental validation on the promotion of CS2 on the

ignition of methane.

Experiments involving CS2 oxidation in combustion processes have been conducted with

tubular-flow (TFR) and jet-stirred (JSR) reactors to provide experimental validation for the

proposed mechanism. Online Fourier transform infrared (FTIR) spectroscopy served to

identify and quantitate the product species and to obtain the detailed species conversion profiles

during the combustion process. Low ignition temperature of CS2 of 860 K in TFR and 710 K

JSR, with residence time at 0.3 s under ambient pressure, implies the extreme flammability of

CS2. The presence of moisture exhibits no effect on the oxidation of CS2 for ignition

temperature and species profile at below 1100 K, although moisture converts CO into CO2 at

higher temperature (> 1200 K). Co-oxidation experiments of CH4/CS2/O2 in JSR illustrate the

promotion effect of CS2 on the ignition of methane.

Quantum calculations afford the investigation of primary steps governing the OH-initiated

oxidation of CS2 in the atmosphere. We also propose a comprehensive oxidation mechanism

of CS2 in combustion systems. The thesis suggests the intersystem crossing (ISC) between

triplet and singlet pathways to explain the extreme flammability of CS2, in analogy to the

v

controlling steps operating for other reduced sulfur species. DFT calculations examine the

interplay between the SH radical and C1 - C4 hydrocarbons to demonstrate the distinct

inhibition and promotion effects of H2S/SH on alkanes and alkenes/alkynes in pyrolysis

processes. The findings reported in this study apply to both atmospheric and combustion

systems. Especially, the developed mechanism provides improved understanding of the

oxidation of fossil fuels containing sulfur species.

vi

Table of contents

Statement of originality ............................................................................................................ i

Supervisory statement ............................................................................................................. ii

Acknowledgment .................................................................................................................... iii

Abstract .................................................................................................................................... iv

Table of contents ..................................................................................................................... vi

List of publications ................................................................................................................. xii

Chapter 1. Introduction .......................................................................................................... 1

1.1. Research Background ...................................................................................................... 2

1.2. Project objectives and thesis outline ............................................................................... 3

Reference ................................................................................................................................ 7

Chapter 2. Literature review ................................................................................................ 10

2.1. Introduction ................................................................................................................... 11

2.2. Oxidation of H2S ........................................................................................................... 14

2.2.1. Atmospheric oxidation of H2S ................................................................................ 14

2.2.2 Combustion mechanism of H2S ............................................................................... 20

2.3. Oxidation of CS2 ........................................................................................................... 27

2.3.1. Atmospheric oxidation of CS2 ................................................................................ 28

2.3.2. Combustion mechanism of CS2 .............................................................................. 32

2.3.3. Production of S atom in excited state in photolysis of reduced sulfur species ....... 40

vii

2.4. Interaction between reduced sulfur species and hydrocarbons ..................................... 43

2.4.1. Influence of H2S/SH on the combustion process of hydrocarbons ........................ 44

2.4.2. Impact of SO2 on ethylene pyrolysis ...................................................................... 46

2.5. Conclusion ..................................................................................................................... 48

Reference .............................................................................................................................. 52

Chapter 3. Methodology ........................................................................................................ 66

3.1. Experimental methodology ........................................................................................... 67

3.1.1. Gas feeding system ................................................................................................. 67

3.1.2. Tubular-flow reactor (TFR) .................................................................................... 68

3.1.3. Jet-stirred reactor (JSR) .......................................................................................... 69

3.1.4. Online analytical technique .................................................................................... 71

3.2. Computational methodology ......................................................................................... 72

3.2.1. Quantum chemistry calculation .............................................................................. 72

3.2.2. Kinetic modelling ................................................................................................... 77

Reference .............................................................................................................................. 79

Chapter 4. Atmospheric oxidation of carbon disulfide (CS2) ............................................ 80

Abstract ................................................................................................................................ 81

4.1. Introduction ................................................................................................................... 82

4.2. Computational Methodology ........................................................................................ 84

4.3. Results and Discussion .................................................................................................. 85

4.3.1. CS2 + OH ................................................................................................................ 85

viii

4.3.2. S-adduct SCS(OH) + O2 ......................................................................................... 92

4.3.3. C-adduct SC(OH)S + O2 ........................................................................................ 96

4.4. Conclusion ..................................................................................................................... 98

Acknowledgement ................................................................................................................ 98

Reference .............................................................................................................................. 99

Chapter 5. Inhibition and promotion of pyrolysis by hydrogen sulfide (H2S) and

sulfanyl radical (SH) ............................................................................................................ 102

Abstract .............................................................................................................................. 103

5.1. Introduction ................................................................................................................. 104

5.2. Computational Details ................................................................................................. 106

5.3. Results and discussion ................................................................................................. 107

5.3.1. Overview of labile H abstraction sites in C1-C4 hydrocarbons ............................ 107

5.3.2. H abstraction from alkanes ................................................................................... 111

5.3.3. H abstraction from alkenes and alkynes ............................................................... 118

5.3.4. Validation of CBS-QB3 calculations ................................................................... 121

5.3.5. Relationship between BDH of abstracted C‒H bond and activation enthalpy of SH

+ hydrocarbon reaction ................................................................................................... 123

5.3.6. Activity of SH radical compared with those of OH, NH2 and HO2 ..................... 125

5.4. Conclusion ................................................................................................................... 126

Appendix ............................................................................................................................ 127

Acknowledgement .............................................................................................................. 127

Reference ............................................................................................................................ 128

ix

Chapter 6. Flammability of CS2 and other reduced sulfur species ................................. 132

Abstract .............................................................................................................................. 133

6.1. Introduction ................................................................................................................. 134

6.2. Methodology ............................................................................................................... 137

6.2.1. Experiments with tubular flow reactor ................................................................. 137

6.2.2. Kinetic modelling ................................................................................................. 140

6.3. Results and discussion ................................................................................................. 141

6.3.1. Experimental results from tubular flow reactor .................................................... 141

6.3.2. Sensitivity analysis of CS2 consumption .............................................................. 146

6.3.3. CS2/O2 subset ........................................................................................................ 147

6.3.4. S/O2 subset ............................................................................................................ 153

6.3.5. Updated mechanism and modelling validation of our experimental results ........ 155

6.4. Conclusion ................................................................................................................... 156


Reference ............................................................................................................................ 158

Chapter 7. Combustion chemistry of carbon disulfide (CS2) .......................................... 164

Abstract .............................................................................................................................. 165

7.1. Introduction ................................................................................................................. 166

7.2. Experimental ............................................................................................................... 168

7.3. Kinetic Modelling ....................................................................................................... 171

7.4. Results and Discussion ................................................................................................ 176

x

7.4.1 Experimental results with JSR ............................................................................... 176

7.4.2. Kinetic modelling of dry oxidation of CS2 and revision of COS/O2 subset ......... 180

7.4.3. The influence of moisture on CS2 oxidation ........................................................ 187

7.5. Conclusion ................................................................................................................... 191


Appendix ............................................................................................................................ 192

Reference:........................................................................................................................... 193

Chapter 8. Enhanced oxidation of CH4 in presence of CS2 ............................................. 198

Abstract .............................................................................................................................. 199

8.1 Introduction .................................................................................................................. 200

8.2 Methodology ................................................................................................................ 202

8.2.1 Experimental set-up ............................................................................................... 202

8.2.2 Kinetic modelling .................................................................................................. 204

8.3 Results and discussion .................................................................................................. 204

8.3.1 Sole oxidation of CS2/O2 and CH4/O2 ................................................................... 204

8.3.2 Co-oxidation of CS2/CH4/O2 ................................................................................. 206

8.3.3 Kinetic modelling of the co-oxidation of CS2/CH4/O2 .......................................... 209

8.4 Conclusion .................................................................................................................... 213

Reference ............................................................................................................................ 214

Chapter 9. Conclusion and recommendation .................................................................... 217

9.1. Conclusion ................................................................................................................... 218

xi

9.2 Recommendation .......................................................................................................... 220

Appendix A for chapter 5 .................................................................................................... 223

Appendix B for chapter 7 .................................................................................................... 243

Appendix C for risk assessment.......................................................................................... 274

xii

List of publications

Journal articles

1. Z. Zeng, M. Altarawneh, B. Z. Dlugogorski, Atmospheric oxidation of carbon

disulfide (CS2), Chem. Phys. Lett., 669 (2017) 43-48.

2. Z. Zeng, M. Altarawneh, I. Oluwoye, P. Glarborg, B. Z. Dlugogorski, Inhibition and

promotion of pyrolysis by hydrogen sulfide (H2S) and sulfanyl radical (SH), J. Phys.

Chem. A, 120 (2016) 8941-8948.

3. Z. Zeng, B. Z. Dlugogorski, M. Altarawneh, Flammability of CS2 and other reduced

sulfur species, Fire Saf. J., in press, DOI: https://doi.org/10.1016/j.firesaf.2017.03.073.

Conference papers

1. Z. Zeng, M. Altarawneh, B. Z. Dlugogorski, “Reactions of SH radical with C1-C4

Hydrocarbons.” Proceedings of the Australian Combustion Symposium, 2015,

Melbourne, Australia

2. Z. Zeng, B. Z. Dlugogorski, M. Altarawneh, “Flammability of CS2 and Other

Reduced Sulfur Species.” 12th International Symposium on Fire Safety Science (the

12th IAFSS Symposium), 2017, Lund, Sweden

3. Z. Zeng, M. Altarawneh, B. Z. Dlugogorski, “Enhanced oxidation of methane by

CS2.” Asia-Pacific Conference on Combustion, 2017, Sydney, Australia.

Chapter 1. Introduction

1

Chapter 1

Introduction


2

1.1. Research Background

Sulfur bearing compounds represent a group of major impurities in most organic fuels [1, 2].

In combustion, sulfur transforms primarily into oxide species, i.e., SOx, acting as the critical

source for a series of atmospheric air contaminants. The typical sulfur content in coal varies

from 0.4 - 4.0 wt. % [3]. In commercial diesel and gasoline, sulfur is limited to 500 ppm and

30 ppm, respectively [4]. Likewise, pipeline-quality natural gas usually contains sulfur at

levels of around 2,000 ppm [5]. In typical biomass, the content of sulfur varies from 0.02 - 0.3

wt. % [2]. The enormous consumption of fossil fuels in oxidation processes, e.g., in power

plants, engines and stoves, results in significant emission of sulfur oxides, especially in

countries that do not require or enforce the emissions of SOx, limiting the exploitation of fossil

and biomass fuels.

Sulfur dioxide (SO2) constitutes the most notorious sulfur compound emitted from combustion

of fossil fuels. However, under fuel rich combustion conditions or during thermal pyrolysis of

fuels, the product gases may also contain reduces sulfur species, such as hydrogen sulfide

(H2S), carbon disulfide (CS2) and carbonyl sulfide (COS). Gas processing plants and oil

purification and gasification processes produce H2S as a side product. While CS2 and COS

evolve as major by-products in Claus gas-desulfurising process [6, 7], which converts H2S into

elementary solid sulfur, volcanic eruptions and bushfires [8] also contribute to the total

atmospheric budget of CS2 and COS. CS2 exists at levels of parts per trillion (ppt) in the earth’s

troposphere; 15 to 30 ppt in the nonurban troposphere and 100 to 200 ppt in polluted urban

areas [9]. On the other hand, COS stands as the most abundant sulfur compound in the

atmosphere (> 400 ppt) with a relatively long life-time of up to years [10].


3

Existing as a volatile liquid at room temperature (25 oC), CS2 represents a non-polar solvent

which has found numerous applications in laboratories and industry, such as extractive

metallurgy. Likewise, the production of synthetic fibres, rubbers and pesticides requires CS2

and COS as feedstocks. Unfortunately, CS2 displays extreme flammability and high explosion

propensity [11], properties responsible for fires in laboratories and chemical warehouses [12,

13].

Apart from H2S [14], the oxidation mechanism and combustion properties of CS2 and COS

remain poorly understood. Lack of appropriate reaction mechanisms prevent detailed

modelling of important industrial plants, such as the Claus process or separating sulfur

impurities in the feedstocks in oil refineries [15]. To gain insight into the extreme flammability

of CS2 and the role of sulfur compounds in oxidation of fuels, one needs not only the

mechanisms of oxidation of CS2 and COS, but also the mechanisms describing the interaction

between sulfur and hydrocarbon species.

1.2. Project objectives and thesis outline

The general objective of this study is to construct a comprehensive oxidation mechanism for

CS2 and to investigate the reactions between sulfur species and hydrocarbons. Specifically,

this thesis intends to:

1. Perform accurate quantum chemistry calculations to improve the existing mechanism

of atmospheric oxidation of CS2


4

2. Collect precise measurements of CS2 oxidation under combustion condition using

tubular-flow (TFR) and jet-stirred (JSR) reactors

3. Investigate the extreme flammability and low ignition temperature of CS2 with quantum

chemistry calculations and propose further enhancements to the oxidation mechanism

for CS2

4. Validate the proposed mechanism against the experimental results from TFR and JSR

through kinetic modelling

5. Study the interaction between SH radical and C1 – C4 hydrocarbons with density

function theory (DFT) calculations, and

6. Explore the effect of CS2 addition on the ignition of methane with careful JSR

experiments.

To address the above aims, the thesis has the following structure:

Chapter 2 reviews the literature pertinent to experiments and reaction kinetics of reduced sulfur

species. We highlight the occurrence of intersystem crossing (ISC) in H2S and S/SO oxidation

processes. ISC is a transition between two electronic states with different spin multiplicities

but same energy level. In the case of sulfur chemistry, the oxidation reaction will start with

triplet oxygen molecule and transit to singlet reaction pathway through the cross-over point on

the potential energy surfaces, offering an alternative reaction corridor with much lower

activation barriers. For CS2 oxidation, the high activation energy of triplet oxidation pathway

fails to explain the low ignition temperature for the experiments conducted with the tubular-

flow reactor, later prompting us to introduce the ISC to the CS2 combustion mechanism.

Photolysis of reduced sulfur species (H2S, CS2 and COS) also leads to the formation of singlet

sulfur atom (1S). This chapter also examines the interaction between sulfur species and


5

hydrocarbons, summarising 1) influence of H2S/SH on hydrocarbons combustion process and

2) impact of SO2 on the pyrolysis of ethylene.

Chapter 3 covers the experimental and theoretical methodologies applied in this project. It

depicts the experimental set-up constructed in this work to study the oxidation of CS2 and its

interaction with hydrocarbons, including the reactant feeding system, details of TFR and JSR

and online analytical measurements. This chapter also comprises a description of the quantum

chemistry calculations and kinetic modelling, including the software used in this thesis:

Gaussian 09, MESMER 3.0, ChemRate, KiSTheIP and Chemkin Pro.

Chapter 4 investigates primary steps governing the OH-initiated atmospheric oxidation of CS2

with theoretical calculations. In the absence of surface effects, we find the overall reaction OH

+ CS2 → COS + SH too slow to account for the formation of the reported experimental

products. The S-adduct represents the most plausible product formed by addition of OH to

CS2. The adduct then undergoes a bimolecular reaction with atmospheric O2 yielding COS,

OH and SO. The kinetic analysis developed in this study explains the atmospheric fate of

reduced sulfur species, an important group of compounds in the global S cycle.

Chapter 5 resolves the interaction of sulfanyl radical (SH) with aliphatic (C1 - C4)

hydrocarbons, using CBS-QB3 based quantum chemistry calculations. Our findings

demonstrate that, the documented inhibition effect of hydrogen sulfide (H2S) on pyrolysis of

alkanes does not apply to alkenes and alkynes. During interaction with hydrocarbons, the

inhibitive effect of H2S and promoting interaction of SH radical depend on the reversibility of

the H abstraction processes. In the case of methane, we conclude that, the reactivity of SH


6

radicals towards abstracting H atoms exceeds that of HO2 but falls below those of OH and NH2

radicals.

Chapter 6 discusses the experimental results of CS2 oxidation using a tubular-flow reactor

under dry conditions for an oxygen-fuel equivalence ratio of 0.7, 1.0 and 1.2 and for a

temperature range of 700 to 1200 K. Compared with moist oxidation, the reaction forms CO

as a final product, as we detected no CO2 in the exhaust gases. A sensitivity analysis confirmed

a significant influence of the CS2/O2 and S/O2 sub-mechanisms on the oxidation process.

Subsequent theoretical calculations identify the intersystem crossing between triplet and singlet

pathways that lowers the activation energy of the key elementary reactions, as the reason for

the extreme flammability of CS2. This chapter also presents the analysis of the nature of the

intersystem crossing and its impact on the low ignition temperature of CS2, and updates the

CS2/O2 and S/O2 reaction subsets for the dry oxidation mechanism of CS2. The results of the

kinetic modelling and the experimental measurements demonstrate good agreement.

Chapter 7 introduces an experimental study of dry and moist oxidation of CS2 in JSR and

develops a comprehensive oxidation mechanism with updated COS/O2 subset. The dry and

moist oxidation of CS2 exhibit the same conversion profile for each species in experiments

performed in the temperature range of 650 – 1100 K under the atmospheric pressure. In

comparison, kinetic modelling with the moist oxidation mechanism proposed in this work

predicts the conversion of CO to CO2 at temperatures in excess of 1200 K. The fast conversion

of COS in experiments, in which this species arises as an intermediate, forces us to study in

detail the consumption channels for COS. We propose the intersystem crossing process to

occur in COS/O2 subset, in analogy to the ISC appearing for other species in the CS2/S/O2


7

system, to reduce the reaction barrier. Good agreement emerges between measured and

modelled onset temperature and CS2 consumption, confirming the robustness of the model.

Chapter 8 explores the impact of added CS2 in the oxidation of CH4. Co-oxidation experiments

of CH4/CS2/O2 in JSR illustrate the promotion effect of CS2 on the ignition of methane in

natural gas. However, the presence of CH4 significantly elevates the oxidation temperature of

CS2, indicating the inhibition effect of CH4 on CS2 oxidation.

Chapter 9 presents the concluding remarks of the thesis and provides suggestions for future

studies on chemistry of reduced sulfur species.

Reference

[1] C.F. Cullis, M.F.R. Mulcahy, The kinetics of combustion of gaseous sulphur compounds,

Combust. Flame 18 (1972) 225-292.

[2] A. Williams, J.M. Jones, L. Ma, M. Pourkashanian, Pollutants from the combustion of solid

biomass fuels, Prog. Energy Combust. Sci. 38 (2012) 113-137.

[3] G.P. Huffman, S. Mitra, F.E. Huggins, N. Shah, S. Vaidya, F. Lu, Quantitative analysis of

all major forms of sulfur in coal by x-ray absorption fine structure spectroscopy, Energy &

Fuels 5 (1991) 574-581.

[4] Department of Environment and Energy, Fuel Standard Determination 2001, Government

of Australia, 2001.


8

[5] P. Jaramillo, W.M. Griffin, H.S. Matthews, Comparative life-cycle air emissions of coal,

domestic natural gas, LNG, and SNG for electricity generation, Environ. Sci. Technol. 41

(2007) 6290-6296.

[6] K. Karan, A.K. Mehrotra, L.A. Behie, A high-temperature experimental and modeling

study of homogeneous gas-phase COS reactions applied to Claus plants, Chem. Eng. Sci. 54

(1999) 2999-3006.

[7] K. Karan, L.A. Behie, CS2 formation in the Claus reaction furnace: a kinetic study of

methane−sulfur and methane−hydrogen sulfide reactions, Ind. Eng. Chem. Res., 43 (2004)

3304-3313.

[8] P. Warneck, The biogeochemical cycling of sulfur and nitrogen in the remote atmosphere,

EOS, Trans. Am. Geophys. Union, 68 (1987) 240-240.

[9] V. Gardiner, Global Tropospheric Chemistry—A Plan for Action, National Academy Press,

1985. 110-111.

[10] M.P. Barkley, P.I. Palmer, C.D. Boone, P.F. Bernath, P. Suntharalingam, Global

distributions of carbonyl sulfide in the upper troposphere and stratosphere, Geophys. Res. Lett.

35 (2008) L14810.

[11] F.R. Taylor, A.L. Myerson, First limit induction time studies of CS2−O2 explosions, Symp.

(Int.) Combust. [Proc.] 7 (1958) 72-79.

[12] American Industrial Hygiene Association, Carbon Disulfide (CS2) Fire.

https://www.aiha.org/get-involved/VolunteerGroups/LabHSCommittee/Pages/Carbon-

Disulfide-Fire.aspx (accessed 05/05 2017).

[13] G.M. Bodner, Lecture demonstration accidents from which we can learn, J. Chem. Educ.

62 (1985) 1105.

[14] F.G. Cerru, A. Kronenburg, R.P. Lindstedt, Systematically reduced chemical mechanisms

for sulfur oxidation and pyrolysis, Combust. Flame 146 (2006) 437-455.


9

[15] F. Viteri, M. Abián, Á. Millera, R. Bilbao, M.U. Alzueta, Ethylene–SO2 interaction under

sooting conditions: PAH formation, Fuel 184 (2016) 966-972.

Chapter 2. Literature review

10

Chapter 2

Literature review


11

2.1. Introduction

Sulfur stands as an essential impurity in fossil fuels, biofuels and municipal wastes [1, 2].

Taking up to several percentage by weight in fuels, sulfur gives rise to industrial problems on

a large scale, especially as an atmospheric pollutant [3]. In nature, sulfur occurs in the

pure element form of sulfide and sulfate in minerals. Once released into air through oxidation,

the mainly gaseous product - sulfur oxide (SOx) is in the form of sulfur dioxide (SO2) [4],

which is an important greenhouse gas as well as the source of acid rain. Anthropogenic

emissions of sulfur into the troposphere peaked during the year 1972 at about 131 million

tonnes [5]. By 2011, due to the world-wide efforts to control air pollution, the emission of

sulfur species has been reduced to 106 million tonnes [6]. More than 85 % of sulfur compounds

originate from fuel combustions in power plants or other devices [7]. As a result, sulfur

chemistry draws a lot of attention among researchers in both energy and environmental

engineering.

Most of the current research on sulfur combustion focuses on the industrial thermal conversion,

oxidation of H2S and emission of SO2 as air pollutant. H2S emerges as the principal sulfur

carrier, in the gasification process of coal [8]. In oil industries, sulfur compounds such as thiols,

thiophenes, sulphides and disulfides (0 – 3 wt. %), are converted to H2S, before they could

poison the cracking catalysts for hydrocarbons [9]. H2S is also commonly present in natural

gas [10]. To remove H2S, amine treatment is applied to crude oil [11] while cation containing

materials, such as limestone (Ca2+) and Magnesium oxide (Mg2+), are sprayed into coals [12].

The multi-step Claus process can convert sulfur into the solid state from gaseous H2S in natural

gas and other exhausts containing H2S, derived from refining crude oil and other industrial

processes [13]. The oxidation of H2S releases SO2, known as the most notorious sulfur


12

compound emitted from fuel combustion [14]. However, a comprehensive oxidation

mechanism for H2S still needs further studies, with key elementary reactions plagued by

significant uncertainties [15].

Apart from H2S, other less-known reduced sulfur species – carbon disulfide (CS2) and carbonyl

sulfide (COS) – are also produced during combustion or thermal pyrolysis of fuels. H2S, SO2,

CS2 and COS are major sulfur-containing species produced during coal pyrolysis [16]. In the

Claus process, CS2 and COS contribute up to 20 % of the total sulfur emissions in exhaust [17,

18]. The formation of CS2 and COS is confirmed in the oxidation of methane seeded with H2S

[19, 20], which is prevalent in natural gas as an impurity. The presence of SO2 in ethylene

pyrolysis also leads to the formation of CS2 [21]. However, studies on CS2 and COS oxidation

are scarce, leading to uncertainties in modelling combustion processes in the real-wold

scenario.

Pure CS2 exists as a colourless liquid, with a low boiling point at 46.3°C under atmospheric

pressure. As an excellent non-polar solvent, CS2 has found numerous applications in both

laboratories and industry. Additionally, the production of synthetic fibres, rubbers and

pesticides requires CS2 as feedstock. In industry, CS2 is manufactured with methane and solid

sulfur, in the presence of silica gel or alumina catalyst under 600 °C [22]:

2 CH4 + S8 catalyst, 600°C→ 2 CS2 + 4 H2S

However, CS2 is characterized by extreme flammability [23] and high explosion propensity

[24], which has led to several fire accidents in laboratories [25, 26] and chemical warehouses

[27].


13

Natural production of CS2 and COS occurs from volcanic eruptions and bushfires [28], while

anthropogenic emission contributes to most of the total atmospheric content of CS2 and COS

[29]. CS2 exists at levels of parts per trillion (ppt) in the earth’s troposphere; 15 to 30 ppt in

the non-urban troposphere and 100 to 200 ppt in polluted urban areas [30]. On the other hand,

COS stands as the most abundant sulfur compound in the atmosphere (> 400 ppt) [31].

Despite the important role of reduced sulfur species in the environmental and energy fields,

studies on combustion chemistry of CS2 and COS are still limited. This study aims to review

the literature related to experiments and oxidation kinetics of reduced sulfur species (H2S, CS2

and COS), and their interaction with hydrocarbons in combustion processes. Firstly, we

examine the pre-established oxidation processes for H2S, which contains S/SO sub-mechanism

SH radical as the key intermediate. The occurrence of intersystem crossing (ISC) in S/SO and

H2S combustion processes is highlighted. Next, we survey the oxidation of CS2 and COS. In

the atmosphere, CS2 will convert itself to COS with the aid of OH radical. COS represents the

most stable and most abundant of sulfur carrier in the atmosphere. For CS2 oxidation, the high

activation energy of triplet oxidation pathway fails to explain the low ignition temperature for

the experiments conducted with the tubular-flow reactor. The ISC is highlighted to occur in

the CS2 combustion mechanism. COS is produced as an intermediate in CS2 oxidation process,

i.e., COS/O2 mechanism is involved in CS2 oxidation as a subset. We also discuss the

production of S atom in both ground state (3S) and excited state (1S) in photolysis of reduced

sulfur species. Finally, the interaction between sulfur species and hydrocarbons has been

investigated for 1) Influence of H2S/SH on combustion process of hydrocarbons and 2) impact

of SO2 on ethylene pyrolysis. Future research directions for sulfur chemistry will be suggested

in the conclusion section, for both experimental and theoretical aspects.


14

2.2. Oxidation of H2S

As the most significant sulfur carrier in natural gas, oxidation and pyrolysis of H2S has been a

subject of many studies focusing on its formation in coal pyrolysis [32], conversion in Claus

process [13] and application as potential hydrogen source [33, 34]. As a flammable gas, the

oxidation of H2S has been investigated with batch reactor to reveal its explosion limit [35, 36],

and shock wave tube to decide its ignition delay time [37]. With the aid of a mass spectrometer,

SO, SO2 and SO3 are sampled from the flat flame of H2S [38].

The recent study by Zhou et al. proposes a comprehensive oxidation mechanism for H2S with

tubular flow reactor [14] and accurate theoretical calculations [39, 40]. Song et al. extended

both the experimental and kinetic modelling of H2S oxidation in plug flow reactor at elevated

pressure up to 100 bar [15]. Owing to the uncertainties in reactions related to SH radical,

further work is necessary to improve the mechanism of H2S oxidation.

In this section, H2S oxidation is reviewed under atmospheric and combustion processes. We

also make recommendations for future work on H2S oxidation in each part.

2.2.1. Atmospheric oxidation of H2S

Reduced sulfur species such as H2S, CS2 and COS represent major atmospheric S-carriers and

assume an important role in the global sulfur cycle. It is generally believed that the majority


15

of H2S is contributed by natural processes such as volcanic eruptions, bush fires, life-cycles in

marshes and sea [28, 41]. However, human activities also result in a significant emission of

H2S, especially in the energy industry [42]. In earth’s troposphere, H2S varies in the level of

20 – 60 ppt [43].

Direct interaction between H2S and O2 proceeds slowly under atmospheric thermal conditions.

While no direct experimental measurement has been found, a quantum chemistry calculation

was conducted by Montoya et al. [44] to demonstrate a high activation energy at 159.8 kJ/mol.

The reaction rate for R2-1 is expressed as k1 = 4.6 × 10-19 × T2.76 × exp (-159.8/RT)

[cm3/(molecule∙s)] in Arrhenius equation fitted from 300 to 2000 K. As extrapolated in this

work, the rate constant for H2S + O2 corresponds to 2.7 × 10-40 cm3/(molecule∙s), which is too

slow to account for the atmospheric conversion of H2S (even when the atmospheric oxygen

concentration is considered).

H2S + O2 → HO2 + SH R2-1

With analytical techniques applied to atmospheric chemistry, researchers have investigated the

oxidation agent of H2S in the atmosphere among atomic oxygen O [45], ozone O3 [46], atomic

hydrogen H and hydroxyl radical OH. The reaction between H2S and O has been studied with

the aid of electron spin resonance (ESR) applied detect and measure O concentration in a flow

reactor [45]. UV-Vis absorption of O3 and resonance fluorescence of H facilitate researchers

in measuring the reaction rate of H2S + O3 [46] and H2S + H [47]. The reaction between OH

and H2S has been measured in discharge-flow reactor with laser induced fluorescence LIF, to

quantitate the concentration of OH radical [48].


16

From Table 2.1, which contrasts the rate constant of H2S oxidized by O2, O, O3, H and OH in

the atmosphere, it can be observed that the reaction between OH radical and H2S contributes

to the most important consumption channel for atmospheric H2S. Moreover, the very low

concentrations of O and H compared to OH in the atmosphere [49] clearly weakens the

influence of R2-2 and R2-4.

Table 2.1. Atmospheric oxidation of H2S, reaction rate k (cm3/(molecule∙s)) is under

atmospheric condition at 298 K, 1 atm.

Reaction Method k Source

R2-1 H2S + O2 → HO2 + SH Theoretical

calculations

2.7 × 10-40 [44]

R2-2 H2S + O → OH + SH Experimental

measurement

2.3 × 10-14 [45]

R2-3 H2S + O3 → HO2 + SH Experimental

measurement

4.0 × 10-16 [46]

R2-4 H2S + H → H2 + SH Experimental

measurement

9.6 × 10-13 [47]

R2-5 H2S + OH → H2O + SH Experimental

measurement

3.6 × 10-12 [48]

By locating OH to be the most potent oxidisers for H2S in the atmosphere, the reaction pathway

of R2-5 has been studied intensively with quantum chemistry calculations by Mousavipour et

al. [50]. Potential energy surface is explored at the MP2/6-311++G(d,p) level of theory [51].

Rate constants were calculated using transition state theory (TST) [52] with one-dimensional

correction for tunneling effect [53], resulting in an activation energy at 4.2 kJ/mol fitted from

300 – 3000 K. The computed reaction rate amounts to 2.65 × 10-12 cm3/(molecule∙s) at 298 K,

which is in good agreement with the experimental analogous value at 3.6 × 10-12

cm3/(molecule∙s) [48], confirming the accuracy of the theoretical calculation.


17

Figure 2.1 demonstrates the reaction pathway and structure of the transition state structure

calculated by Mousavipour et al. [50]. The highly reactive OH radical will extract one H from

H2S to form H2O and SH radical. This reaction can be understood as H exchange between SH

and OH radical. Since the bond dissociation enthalpy of H2O (498.8 kJ/mol) is much higher

than that of H2S (380.1 kJ/mol) at 298 K [54], the OH radical is far more competitive than the

SH radical, to acquire H.

Figure 2.1. Calculated reaction pathway and transition state structure for H2S + OH → H2O +

SH, plotted from calculation by Mousavipour et al. [50]. All enthalpy values are with reference

to initial reactants H2S + OH, in kJ/mol at 0 K. All distance values (around chemical bonds)

are in Å = 10-10 m.

Further conversion of SH radicals, which involves O2, would proceed through R2-6: SH + O2

→ SO + OH. However, the reaction rate of R2-6 is measured to be less than 1.0 × 10-17

cm3/(molecule∙s) [55] using resonance fluorescence technique, which is too slow to account

for the consumption of SH in the atmosphere. O3 and NO2 have been found to be the major

agents for SH conversion at a reasonable reaction rate, to produce HSO [56, 57] as products.


18

As shown in Table 2.2, NO2 acts as the most powerful scavenger for SH in the atmosphere. In

spite of the important role of NO2 in atmospheric sulfur cycle, no quantum chemistry

calculation has been conducted to reveal the reaction pathway of R2-8.

Resende and Ornellas [58] applied theoretical calculations to explore the consumption of SH

in the atmosphere for reaction R2-7: SH + O3 → HSO + O2. However, by conducting

calculations in ground state with 2SH and 3O3, the computed activation energy attains a value

of 26.5 kJ/mol. This energy barrier leads to a reaction rate for R2-7 at 2.3 × 10-17

cm3/(molecule∙s) for atmospheric condition at 1 atm and 298 K, which is too slow compared

to experimental measurement at 4.3 × 10-12 cm3/(molecule∙s). Since no spin forbidden channels

were considered in this work, we suggest re-investigation of the reaction, proceeding through

the excited state of reactants (1O3) or transition state. The intersystem crossing between triplet

and singlet pathway may account for the fast conversion of SH as measured in R2-7.

Table 2.2. Atmospheric conversion of SH, reaction rate k is under atmospheric condition at

298 K, 101 kPa.

Reaction Method k (cm3/(molecule∙s)) Source

R2-6 SH + O2 → SO + OH Experimental

measurement

< 1.0 × 10-17 [55]

R2-7 SH + O3 → HSO + O2 Experimental

measurement

4.3 × 10-12 [56]

R2-8 SH + NO2 → HSO + NO Experimental

measurement

6.5 × 10-11 [57]

The reaction rate of R2-9: HSO + O2 → HO2 + SO is measured to be less than 2.0 × 10-17

cm3/(molecule∙s) [59] under atmospheric condition. While the interaction between HSO and

O3 in the atmosphere (R2-10: HSO + O3 → HSO2 + O2) is also too slow to afford the


19

consumption of HSO in the atmosphere with a measured reaction rate at 5.0 × 10-14

cm3/(molecule∙s) [60]. Reaction with NO2 (R2-11: HSO + NO2 → HSO2 + NO) has been

found to be the main channel of HSO conversion, with a reaction rate measured to be 9.6 ×

10−12 cm3/(molecule∙s) in the atmosphere [59]. To the best of our knowledge, literature

presents no account of the reaction between HSO and OH. In this regard, the OH radical would

extract the weakly bonded H atom from S-H bond in HSO to form H2O and SO.

Table 2.3. Atmospheric conversion of HSO and HSO2, reaction rate k is under atmospheric

condition at 298 K, 101 kPa.

Reaction Method k (cm3/(molecule∙s)) Source

R2-9 HSO + O2 → HO2 + SO Experimental

measurement

< 2.0 × 10-17 [59]

R2-10 HSO + O3 → HSO2 + O2 Experimental

measurement

5.0 × 10-14 [60]

R2-11 HSO + NO2 → HSO2 + NO Experimental

measurement

9.6 × 10−12 [59]

R2-12 HSO2 + O2 → SO2 + HO2 Experimental

measurement

3.0 × 10−13 [59]

Subsequent conversion of HSO2 is measured by Lovejoy et al. [59] to proceed with O2 with a

reaction rate at 3.0 × 10−13 cm3/(molecule∙s). However, it must be noted that the methods used

by Lovejoy et al. to obtain the rate coefficient was indirect, which may lead to substantial error

due to the uncertainty of HO2 released from parallel reactions [61]. Table 2.3 lists the reaction

rate for HSO and HSO2 consumption in the atmosphere.

It is worthwhile mentioning that no theoretical studies on the atmospheric conversion of HSO

and HSO2 have been found. Here, we would recommend further studies on atmospheric

reaction involving HSO and HSO2 with computational chemistry, to reveal the detailed reaction


20

pathways. The proposed reaction HSO + OH may open another corridor for HSO consumption,

which is operating independent of the NO2 species in the atmosphere.

Figure 2.2 presents a summary for H2S oxidation in literature, as well as a recommended

pathway, which is of potential importance. In a nutshell, NO2 is extremely important in the

consumption of H2S in the atmosphere. However, in the absence of NO2, an alternative

pathway is proposed here involving OH radical, which is relatively more prevalent in the

atmosphere. We also highlight the failure of ground state calculation on SH + 3O3 to reproduce

the experimental results. The intersystem crossing between triplet and singlet may assist in

reducing the activation energy for the reaction SH + 3O3. Finally, more theoretical work on

HSO and HSO2 is suggested to broaden our understanding on atmospheric oxidation of H2S.

Figure 2.2. Atmospheric oxidation pathway for H2S and suggested study on alternative

pathway.

2.2.2 Combustion mechanism of H2S

Early research on H2S oxidation focused on its ignition delay time [37], explosion limit [35,

36] as well as the flame structures [20, 38], and the reaction mechanism was based on measured


21

product distributions. The first comprehensive review for the combustion of H2S was

conducted by Cullis and Mulcahy in 1972 [3], based on low-temperature photolysis

experiments and flame studies. The final product of H2S oxidation is identified as SO2 with

SO and S2O performing as important intermediates. The proposed consumption channels

involve intermediates such as SH, SO, SO3, S, S2, OH and HO2. However, due to the lack of

reliable kinetic parameters for reactions and thermodynamic data for these species, the

validation of the mechanism remains a very challenging task. Further study by Gargurevich

[62] summarized some early work on kinetics of sulfur chemistry and also estimated the rate

parameters for the missing steps, to resolve the role of H2S in Claus process. Further studies

by Hughes et al. [63] discussed the uncertainties caused by key steps such as SO + OH and SO2

+ H. By realizing the importance of SH, S and SO radical in the mechanism, Tsuchiya et al.

[64] measured the reaction rate for three key elementary reactions: SH + O2, S + O2 and SO +

O2 with shock wave tube at high temperature range (1000 – 1600 K). A systematic reduced

combustion mechanism for H2S has been proposed by Cerru et al [65, 66], extending the radical

pool to contain HSS, HSSH, HSO and HSO2. However, the author attributes the uncertainties

related to SH, HSS and HSSH reactions to insufficient studies, and highlights the paramount

importance of related reactions on the ignition of H2S oxidation.

Some recent studies by Zhou et al. [67, 68] and Gao et al. [39] have combined experimental

work and high level theoretical calculations to resolve the role of H2/S2 system (including SH,

HSS and HSSH) in H2S oxidation, in which they suggest a crossing process from triplet ground

state to singlet excited state for transition structure. The updated mechanism results in a

satisfactory agreement with further experimental validation, in a vertical tubular flow reactor

at 950 – 1150 K under atmospheric pressure [14]. Song et al. further expanded the experiments

with flow reactor under high pressure range between 30 and 99 atm (atmospheric pressure),


22

and interpreted the results with the mechanism of Zhou et al. [14] with certain modification.

It is pointed out that the branch ratio of the reaction of SH + SH to give H2S + S (chain

propagation) or HSSH (chain termination) is the key step to control the oxidation of H2S.

We also intend to point out that the mechanism modelling with plug-flow reactor (PFR) in the

work of Zhou et al. work [14] may not be consistent with their experiments. For tubular flow

reactor, the flow condition can be approximately divided into the laminar [69] and the turbulent

[70] flow regime. As controlled by residence time, the flow condition of inlet gases wouldn’t

reach a turbulent flow with a low flow velocity and Reynold’s number. The species

composition and temperature distribution in laminar flow condition may result in disagreement

with plug flow modelling, in which species and temperature are homogeneous everywhere.

The modelling of PFR is only an approximation to tubular-flow reactor under atmospheric

pressure. This also explains the fact that Song et al. applied a high pressure tubular-flow reactor

to reach turbulent flow condition as modelled with PFR.

In this section, we review the key oxidation process of H2S based on the mechanism of Zhou

et al. [14]. The occurrence of intersystem crossing for selected elementary reactions, especially

in H/S system and S/O system, is the focal point of this part.

The chain initiation of H2S oxidation has been studied by Montoya et al. [44] through

theoretical calculation with G2 method [71]. However, direct H abstraction by O2 in ground

state is featured with a sizable activation enthalpy up to 168.8 kJ/mol at 0 K, as shown in Fig.

2.3. The reaction is also strongly endothermic by 169.5 kJ/mol, which is not favored in

combustion processes. The backward reaction is more prevalent due to a very shallow reverse

reaction barrier. This indicates that SH tends to abstract H from HO2.


23

Figure 2.3. Potential enthalpy surface for H2S + O2 calculated with G2 method at 0 K by

Montoya et al. [44]. All enthalpy values are with reference to initial reactants H2S + 3O2, in

kJ/mol.

However, a crossing between singlet and triplet state is considered by the author to give us a

singlet HOOSH adduct, which is even stabler than the initial triplet reactants H2S + 3O2 by

15.8 kJ/mol. The intersystem crossing (ISC) from ground triplet state to singlet adduct opens

another corridor to form HSO + OH, with a lower activation barrier through the crossing point

(ISC point). However, the author has proposed the ISC process without locating the crossing

point. Here, we would recommend calculation of the detailed potential energy surface (PES)

for both triplet and singlet pathway separately. Then by overlapping the two PESs, we could

approximately fix the crossing point.

H2S + O2 → HSO + OH R2-13

The consumption of HSO has not been investigated. The sole estimation from Zhou et al. [14]

suggests an activation energy at 27.6 kJ/mol to produce SO2 and OH. Here, we would

recommend a theoretical study on the formation of a four-membered ring of HOOSO, where

the O-O bond breaks to give OH and SO2. The produced OH radical extracts one H from H2S


24

readily, through a trivial activation energy at 4.2 kJ/mol, as calculated by Mousavipour et al.

[50]. The quantum calculation of R2-2 has been discussed in the atmospheric part to give SH

as an important chain carrier.

HSO + O2 → SO2 + OH R2-14

H2S + OH → SH + H2O R2-2

Direct measurement of SH + O2 has been conducted by Tsuchiya et al. [64] with shock wave

tube. The activation energy is fitted at 75 kJ/mol in temperature range from 1400 – 1700 K.

However, the theoretical calculation by Zhou et al. [72] at ground state (triplet pathway)

predicted the formation of SO + OH and HSO + O via activation barriers at 81 kJ/mol and 89

kJ/mol at 0 K, respectively. The production of HSO + O is unlikely to proceed with an

endothermicity at 89 kJ/mol, which means the reverse reaction is featured with no activation

barrier. However, the theoretical values derived from this work underestimate the reaction rate

by a factor of around 1.5, compared to measured results from 1400 – 1700 K. A faster reaction

channel is needed to reproduce the experimental results. By considering an electronically

excited state HSO2 suggested by Freitas et al. [73], we recommend underpinning the singlet

reaction pathway and a potential crossing point to ground triplet reactions. This may explain

the relatively fast conversion between SH and O2 as measured in experiments.

SH + O2 → OH + SO R2-15

SH + O2 → HSO + O R2-16

Due to the low reactivity between SH and O2, the recombination of SH + SH is considered to

act as a major chain propagation step. Theoretical calculations were made by Zhou et al. [40,


25

68] with consideration of ISC. As depicted in Fig. 2.4, the triplet reaction pathway is featured

with an activation barrier at 23.7 kJ/mol at 0 K. SH abstracts H from another SH to produce

H2S and S atom. An alternative path is provided through singlet pathway. Direct addition of

SH gives singlet 1HSSH, which further converts to singlet 1H2SS by H migration between S

atoms. Subsequent fission of S-S bond and intersystem crossing from singlet to triplet surface

result in 1H2S and 3S in ground state as products. Further experimental study by Gao et al. [39]

confirms the ISC channel affords the predominant reaction pathway under low temperature

range < 800 K, while triplet pathway would dominate the reaction above 1000 K.

Compared to R2-15, the recombination of SH proceeds through singlet pathway plays an

important role in chain propagating, to produce S atom in the system.

SH + SH → H2S + S R2-17

Figure 2.4. Potential enthalpy surface for SH + SH at 0 K calculated by Zhou et al. [68]. All

enthalpy values are with reference to initial reactants 2SH + 2SH, in kJ/mol.

Further interaction between S and O2 produces SO and O, which exhibits a substantial influence

on the H2S oxidation, as the chain branching step. Lu et al. [74] reported results from combined


26

experimental measurements and theoretical calculations. The experiments reveal that the S +

3O2 reaction demonstrates different temperature dependence for low (T < 1000 K, Ea < 3

kJ/mol) and high (T > 1000 K, Ea > 30 kJ/mol) temperature range, as illustrated in Fig. 2.5.

The behaviour of the S + 3O2 reaction is similar to that of the SH + SH reaction discussed

above. Computational work of the same researchers, performed at the G2M level of theory,

has also explored the possible reaction pathways. The reaction proceeding on a singlet surface

represents the only channel that can account for the low-temperature oxidation. At

temperatures higher than 1200 K, the triplet pathway dominates the overall S + 3O2 reaction.

S + O2 → SO + O R2-18

Tsuchiya et al. [64] observed similar behaviour of SO + 3O2, concluding that the activation

energy of this reaction resides 34.0 kJ/mol above the separated reactants within a temperature

range of 1130 – 1640 K. However, for lower temperatures of 250 – 585 K, the activation

energy corresponds to 19.0 kJ/mol, as reported by Garland [75]. This significant discrepancy

likely originates from ISC phenomenon operating in the low temperature window.

SO + O2 → SO2 + O R2-19


27

Figure 2.5. Arrhenius plot for reaction S + 3O2 based on Lu et al.’s [74] experiments. The

dashed lines represent the triplet (black) and singlet (red) pathways suggested in this work.

In Fig. 2.6, the key consumption steps for H2S oxidation reviewed in this study are

demonstrated. We highlight the ISC process in R2-13, R2-17 and R2-18, and highlight the

possible occurrence of ISC in R2-19. Further investigation on reaction pathway for R2-14 is

also recommended on both triplet and singlet surface.

Figure 2.6. Key oxidation step for H2S in combustion process.

2.3. Oxidation of CS2

In this section, the oxidation pathway of CS2 is reviewed under both atmospheric and

combustion processes. COS oxidation is also included as a sub-mechanism in CS2 combustion


28

mechanism. We also reviewed the photolysis study of reduced sulfur species: H2S, CS2 and

COS, which produce S atom in its excited state (singlet 1S) instead of ground state (triplet 3S).

2.3.1. Atmospheric oxidation of CS2

Apart from H2S, reduced sulfur species of CS2 and COS represent major sulfur carriers in the

atmosphere. Both natural [28] and human activities result in CS2 emission into atmosphere.

However, the majority of atmospheric CS2 originates from industrial processes, such as

production of insecticides and man-made fibres. Moreover, CS2 and COS correspond to an

important by-product from the Claus gas de-sulfurising process [18], taking up nearly 20 % of

the total sulfur emission from the Claus process. CS2 exists at levels of parts per trillion in the

Earth’s troposphere; 15 to 30 ppt in the non-urban troposphere and 100 to 200 ppt in polluted

urban areas [30]. COS stands as the most abundant sulfur compound in the atmosphere (> 400

ppt) with a relatively long life-time of up to years [31].

Similar to H2S, direct interaction between CS2 and O2 is featured by a significant activation

energy, which renders the reaction less important in the atmosphere. With the very low

concentrations of O(3P) in the atmosphere [49] and the absence of any decomposition of CS2

near the UV-region [76], OH radicals constitute the sole initial atmospheric oxidiser. However,

experimental results on the reaction of CS2 and OH reaction indicated an extremely slow rate

constant:

CS2 + OH → COS + SH R2-20


29

By producing OH with pulsed H2O photolysis technique, Kurylo [77] as well as Atkinson et

al. [78] reported the rate coefficient of R2-20 as 1.9 × 10-13 cm3/(molecule∙s), while Atkinson

and Pitts measured an upper limit for R2-20 to be 7.0 × 10-14 cm3/(molecule∙s) at 298.15 K.

However, photolysis of H2O at 165-185 nm may also dissociate the CS2 molecule, resulting in

an error in the reported rate constant of R2-20 due to loss of CS2. A CS2-filling cell was applied

by Wine et al. to filter radiation that dissociates CS2 and report an even lower rate constant for

R2-20, below 2.0 × 10-15 cm3/(molecule∙s) at 298.15 K [79]. Similar result has been achieved

by Iyer and Rowland who produced OH radicals by a continuous wave (CW) photolysis of

H2O2 at 254 nm [80], avoiding the dissociation energy of C-S bond.

By introducing O2 into the system of CS2 and OH, the consumption of CS2 proceeds with a

much higher reaction rate. Jones et al. [81, 82] measured the rate constant of R2-20 in the

presence of O2 (i.e. CS2 + OH + O2) using photolysis of HONO, reporting a rather rapid rate

constant of 2.0 × 10-12 cm3/ (molecule∙s) at 298.15 K. The promotion effect of O2 was also

confirmed by Barnes et al. using CW photolysis [83], who obtained a very similar rate constant.

Two important steps control the atmospheric oxidation of CS2, namely formation of a CS2OH

and reactions of O2 with this adduct:

CS2 + OH → CS2OH R2-21

CS2OH + O2 → products R2-22

Experiments by Murrells et al. [84] and Lovejoy et al. [85] tested a wide temperature range of

249 – 318 K to confirm the formation of OCS, and SO2 as the sole experimental products.


30

Theoretically, calculations of Lunell et al. [86] at the HF/3-21G(d) level of theory predicted

two distinct isomers of CS2OH, viz., the S-adduct – SCS(OH) and the C-adduct – SC(OH)S.

However, an incorrect structure of the SCS(OH) adduct is considered in Lunell et al.’s work

[86], resulting in a strongly endothermic reaction by 104.2 kJ/mol, with respect to the reaction

of CS2 and OH in R2-21. Subsequent computations of McKee [87] at the modified G1 [88]

level of theory, at 298.15 K, yielded reaction enthalpies of -24.7 kJ/mol and -126.8 kJ/mol,

respectively, for the generation of the S- and C-adducts of CS2OH in R2-21. McKee and Wine

[89] further investigated R2-22 with optimised geometries obtained at the B3LYP/6-311+G(d)

level of theory and constructed a pathway for the formation of COS and HOSO. Figure 2.7

depicts the detailed reaction pathway for the S-adduct to interact with O2. The insertion of O2

on C atom of S-adduct results in an intermediate SCS (OH) O2 with a sizable activation barrier

at 28.6 kJ/mol. Subsequent fissions of C-S and O-O bond yield COS and HOSO as products.

However, the calculated well depth of S-adduct amounts to 26.9 kJ/mol, which underestimates

the stability of SCS(OH) as measured at 45.6 ± 4.2 kJ/mol by Murrells et al. [84].

Figure 2.7. Potential enthalpy surface for CS2 + OH + O2 calculated at 0 K by McKee and

Wine [89]. All enthalpy values are with reference to initial reactants CS2 + OH + O2, in kJ/mol.

The HOSO + O2 reaction has a near zero activation barrier, as calculated by McKee and Wine

[89] The generation of HO2 + SO2 should proceed readily in the atmosphere. It explains the


31

observation of SO2 in experimental measurements. R2-22 is updated below to show the

formation of HOSO, while R2-23 demonstrates its consumption channel in the atmosphere.

CS2OH + O2 → COS + HOSO R2-22

HOSO + O2 → HO2 + SO2 R2-23

While SO2 will be rapidly converted to acid rain by OH and O2 as oxidiser [90-93] in the

atmosphere, COS tends to accumulate as the most abundant sulfur carrier in the earth’s

troposphere (> 400 ppt), with a relatively long life-time of up to years [31].

Figure 2.8. Atmospheric oxidation pathway for CS2.

Figure 2.8 depicts the atmospheric conversion of CS2 with COS and SO2 as products. Although

a comprehensive quantum calculation has been carried out by McKee and Wine [89], the

relative low level of theory may lead to significant error, as compared with experiments.

Further calculation with an adequately high level of theory is desirable, to further improve the

atmospheric oxidation mechanism of CS2.


32

2.3.2. Combustion mechanism of CS2

Earlier work on the combustion of CS2 was reviewed by Cullis and Mulcahy [3] . Myerson

and Taylor extensively studied the ignition of CS2-O2 mixtures as an extremely flammable

chemical, as a representative case for branched-chain reactions [24, 94]. It was demonstrated

that a mixture containing as little as 0.03 % (by mol) CS2 in oxygen ignites at 80 oC and 0.05

atm, resulting in sustained cool flame propagation. Once ignited, a cool flame propagates

through the mixture at 55 oC, whereas the temperature rise in the flame is less than 15 oC. The

most important oxidation products were reported to be CO, SO2, CO2, CS and SO, as well as

COS. Intermediates such as SO, S and O appear much more reactive than COS and CS [95],

with relatively low accumulation.

Studies investigated the effect of various factors on the ignition behaviour of CS2, including

irradiation by ultraviolet light [96], the presence of inert gases [97] and the conditions of

surfaces [24]. Results from these investigations pointed to a profound distinction between cool

and hot flames of CS2 with reference to hydrocarbons [98, 99]. As evidence of the cool flame

combustion of CS2, it was found that the flammability of CS2 increases with the concentration

of oxygen in the combustible CS2 mixture [100]. This behaviour significantly differs from that

observed in combustion of hydrocarbons, where the second ignition limit arises due to a three-

body reaction H + O2 + M → HO2 + M [101].

In CS2 oxidation, the two rapid steps R2-18 and R2-19 produce the reactive O radicals:

S + O2 → SO + O R2-18

SO + O2 → SO2 + O R2-19


33

Various chain inhibitors modify the duration of an induction period from a few seconds to

several minutes [24].

Hanst and Myerson et al. [102] were the first to study the kinetics of explosive combustion of

CS2. They used kinetic absorption spectroscopy to follow the variations in the concentrations

of radical and molecular species during the explosion of CS2. The CS2-O2 mixture, initially

heated to between 190 oC and 300 oC, exploded spontaneously. Strong and continuous

absorptions confirmed the formation of SO and CS with the appearance of SO2 in the early

stages of the explosions [96]. Spectroscopic techniques were also utilised to study explosions

of CS2 initiated by shock waves [103] and flash photolysis [104]. In both cases, CS2 was found

to be oxidized mainly into SO, SO2 and CO. However, experimental measurements on slow

combustion and explosive combustion of CS2 have not contributed much insight into

mechanisms governing its oxidation, as only final products at particular low temperature ranges

were determined.

Reaction pathways contributing to the explosion behaviour of CS2 are summarised in Fig. 2.9.

Central to these mechanisms is the slow build-up of CS via reaction of O atom with CS2 and

its further oxidation into SO/SO2. The fate of produced S atom was assumed to be controlled

by reactions with O atom and CS2, which performs as the initiation step for CS2 explosions. It

is worthwhile to note that the formation of an O atom in this mechanism was attributed to

“unknown” initiation reactions.


34

Figure 2.9. Oxidation pathways of CS2 derived from explosion behaviour of CS2 at early stage

[96]. The initiation step is believed to be CS2 + O, while final products are CO and SO2.

Along the same line of enquiry, a low pressure flame of CS2 was studied by Azatyan et al.

[105] to address the high-temperature combustion mechanism, with the electron spin resonance

(ESR) technique to detect O, CS and SO as intermediates. By analysing the relationship

between O and SO in experiments, the authors proposed the formation of COS as a major

conversion channel between O and SO with the reaction: COS + O → CO + SO.

To confirm the occurrence of these reactions, Azatyan et al. [105] added small concentrations

of COS to the CS2 flame at a temperature interval of 350 oC to 600 oC. A lower accumulation

of O and the maximum yield of SO were observed, validating the influence of COS in the

oxidation of CS2.

Homan et al. [106] deployed the isothermal flow reactor to study combustion of CS2 in an

excess of O2 at 927 oC and 0.4 atm. They confirmed the presence of COS in the oxidation

process of CS2. In order to describe their experimental product profiles, a reduced kinetic

model was constructed including COS as intermediate. The reaction CS2 + O → COS + S

affords the production of COS.


35

CS2 + O → CS + SO R2-24

CS + O → CO + S R2-25

CS + O2 → CO + SO R2-26

COS + O → CO + SO R2-27

CS2 + O → COS + S R2-28

Glarborg and Marshall [107] have recently studied the oxidation mechanism of COS with

kinetic modelling. A detailed oxidation mechanism has been proposed, based on evaluation of

data from literature. For the consumption of COS, the author considered three pathways:

COS → CO + S R2-29

COS + O → CO + SO R2-30

COS + O2 → CO + SO2 R2-31

For the thermal decomposition process (R2-29: COS → CO + S), a high activation energy at

260 kJ/mol fitted from experimental work at high temperature (1900 – 3230 K) [108] makes

it too slow to account for the conversion of COS. The COS interaction with atomic oxygen

(R2-30: COS + O → CO + SO) presents as the major consumption channel for COS. The

authors fitted the rate parameters from the experimental measurement by Homman et al. [16]

around 1270 K, indicating an activation energy at 21.8 kJ/mol. The dominant products are CO

+ SO, while CO2 + S could form at higher temperature. With the absence of atomic O for chain

initiation, the rate parameter of reaction between COS and O2, R2-31 (COS + O2 → CO + SO2),

is estimated with an activation energy at 134.4 kJ/mol.


36

Furthermore, to evaluate their model, the authors have compared it in the oxidation of COS

with analogous experimental data obtained in batch reactors [109], flow reactors [106], and

shock tubes [110]. Although the proposed COS oxidation mechanism captured well the

experimental profiles for major species measured with tubular flow reactor by Homann et al.

[111], a slight correction for the onset of reaction has to be made to match the experiment

initiation. The reaction COS + O has been recognized to be very sensitive to the overall

mechanism and intersystem crossing from triplet to singlet surface would occur with high

probability [112].

The first attempt to construct a detailed kinetic model for CS2 oxidation was undertaken by

Howgate and Barr [113] to model CS2-O2 flame. They include the direct interaction between

CS2 and O2 to account for the chain initiation with an activation energy at 179.5 kJ/mol, as

indicated: CS2 + O2 → CS + SO2. Further investigation by Hardy and Gardiner [114] adopted

this chain initiation step in their mechanism and achieved satisfactory agreement with their

experimental measurement for the ignition delay time of CS2 with shock wave tube.

Subsequent measurement of Saito et al. [115] and Murakami et al. [116] revised the activation

energy of R2-32 to be 134.5 kJ/mol and 130.0 kJ/mol, respectively, to match their experimental

results from shock wave tube.

CS2 + O2 → CS + SO2 R2-32

Very recently, Glarborg et al. [117] formulated a kinetic model for oxidation of CS2. Kinetic

data for key CS2/CS + O2 reactions have been investigated on the basis of ab initio calculations.

For reactions CS2 + O2, the authors investigated the reaction pathway on both triplet and singlet

surface, and concluded the products to be COS + SO, instead of formation of CS + SO2.


37

CS2 + O2 → COS + SO R2-33

Figure 2.10. Potential energy surface for CS2 + O2 calculated by Glarborg et al. [117]. All

energy values are with reference to initial reactants CS2 + O2, in kJ/mol.

As demonstrated in Fig. 2.10, the triplet pathway of CS2 + O2 climbs through the CS2O2 adduct,

which dissociates into COS and SO, through an overall reaction barrier of 223 kJ/mol.

However, the singlet oxidation offers an optional pathway with a much lower activation barrier.

Compared to the triplet CS2O2 structure that features a strong O-O bond, the singlet SOOCS

adduct forms through the dissociative addition of oxygen atoms at C and S atom, with a much

lower energy level. By locating the cross-over between the energy surface of triplet and singlet

pathways, Glarborg et al. estimated the crossing point to reside 145 kJ/mol above the triplet

reactants CS2 + 3O2, which coincides with the experimentally interpreted value in the shock

wave tube system (130.0 – 179.5 kJ/mol) [114-116].

With the formation of SO through R2-33, its interaction with O2 gives O as the most important

chain carrier. The subsequent interactions between CS2 and O proceed through three reaction

channels, resulting in CS + SO (R2-24), COS + S (R2-28) and CO + S2 (R2-34) as products.


38

CS2 + O → CO + S2 R2-34

An experimental study using a fast flow reactor concluded that the formation of CS + SO

dominates the overall process, with only a small generation of COS + S and CO + S2 [118].

However, theoretical calculations on the triplet surface support only the formation of CS + SO,

with a small activation energy at 4.2 kJ/mol. To produce COS + S or CO + S2, the reactions

must overcome higher barriers of 41.3 kJ/mol and 33.2 kJ/mol, respectively. Thus, both

reactions should be negligible, when contrasted with the formation of CS + SO [119]. Here,

we would like to suggest the possible singlet reaction pathways to account for the formation of

COS + S and CO + S2 [120], as detected in experiments.

As calculated by Glarborg et al. [117], CS is consumed by O2 with an activation energy at

117.5 kJ/mol on triplet surface. From shock-tube experiments, Murakami et al. reported the

activation energy of R2-26 as 51.0 ± 61.7 kJ/mol [116]. The significant uncertainty from

experiments and high activation barrier derived by triplet reaction pathway is to be noted; we

would like to suspect the occurrence of ISC for R2-26.

CS + O2 → CO + SO R2-26

In general, the mechanism proposed by Glarborg et al. [117] displays good agreement with

experimental results pertinent to ignition delays and explosion limits of CS2. However, their

mechanism tends to over-predict the concentration of species under low temperature

conditions, especially with respect to measurements from flow reactors [111] and shockwave

tubes [115]. The uncertainty of the chain initiation reaction, CS2 + O2, and competition

between the chain branching reactions, CS + O2, accounted for the discrepancy.


39

More recently, Abián et al. [121] studied the moist oxidation of CS2 in a flow reactor at 1 atm

between 127 – 1127 oC under different oxygen-fuel equivalence ratios (λ) of 0.2, 0.7, 1.0, 2.0

and 20.0. They compared their experimental profiles of consumed CS2 with the corresponding

profiles calculated from the mechanism of Glarborg et al. [117] The modelling results from

the mechanism of Glarborg et al. [117] over-predict the onset of the oxidation temperature for

stoichiometry condition (λ = 1.0) by around 260 K. By adjusting the activation energies of the

initiation reaction (CS2 + O2) and the chain propagation step (CS + O2), to lower values, Abián

et al. reproduced their experimentally-measured concentration of species, as illustrated in Fig.

2.11. However, it is worthwhile noting that the mechanism updated by Abián et al. [121] also

underestimates the ignition of CS2 for around 100 K. The introduction of moisture, lack of

control of residence time at different temperatures and plausible involvement of surface-

mediated reactions on the reactor walls in Abian et al.’s experiments [121], make it difficult to

interpret their experimental measurements and to differentiate plausible initiation by catalytic

surfaces and water molecules. Further experiments on CS2 oxidation for kinetic modelling are

required to develop a more accurate mechanism of CS2 oxidation.

Figure 2.11. Comparison between experimental results for CS2 oxidation and the modelling

results derived from the mechanisms of Glarborg et al.[117] and Abián et al. [121].


40

Figure 2.12. Oxidation pathways of CS2 with formation of COS. The initiation step is

considered to be CS2 + O2, while final products are CO and SO2.

To summarise, Fig. 2.12 demonstrates the key oxidation pathway for CS2, with formation of

COS as an important intermediate. From literature, we highlight the occurrence of ISC in R2-

33 as chain initiation step. More theoretical calculations are recommended on the key

elementary reactions such as R2-26, R2-30 and R2-31, with consideration of singlet reaction

surface. Experimental data for kinetic modelling is also desired to validate and update the

oxidation mechanism for CS2.

2.3.3. Production of S atom in excited state in photolysis of reduced sulfur species

With consideration of the ISC process predicted from theoretical calculation for oxidation of

reduced sulfur species, we review the direct experimental measurements of singlet/triplet

species produced in pertinent studies. In the work of Nan et al. [122], the Doppler-broadening


41

fine spectrometry of S is measured by laser induced fluorescence (LIF) at 147 nm to identify

the ratio of 3S/1S. As photolysis of COS at 222 nm light proceeded in a glass cell, the integrated

intensity of 3S profile increased with time, due to the quenching of 1S produced in reaction.

The production of singlet 1S appears predominant in photolysis of COS, with the produced

3S/1S = 0.05/0.95. While further study on the photolysis of H2S gives combined 3S/1S =

0.87/0.13 as products [123]. The mixed sulfur atoms with distinct electron spin state yielded

from photolysis of reduced sulfur species reinforce the occurrence of ISC in sulfur-related

reactions.

Recently, by coupling mass spectrometry (MS) and photon-ionization techniques with

synchrotron vacuum ultra-violet light as photon source, McGivern et al. [124] and Qi et al.

[125] manage to distinguish the 3S and 1S atom with their ionization energy at 10.36 eV (3S)

and 7.61 (1S), respectively. Compared to the traditional electron ionization applied with MS,

photon ionization is featured with a high resolution at 0.1 eV, which is adequate to recognize

the difference in the energy level of 3S and 1S. A study on CS2 photolysis with 193 nm light

reveals that the combined formation of 3S/1S = 0.75/0.25 was conducted by McGivern et al.

[124]. Figure 2.13 illustrates the energy level of CS2 photolysis reactions.


42

Figure 2.13. The energy level of CS2 photolysis reactions (in kJ/mol) with branch ratio

calculated in this work by Gaussian 09 at CBS-QB3 composite [126].

As demonstrated in Fig. 2.13, we observe that the ground state of CS2 is in singlet (1CS2), while

that of S is in triplet (3S). If the photolysis of CS2 proceeds from ground state to ground state,

an ISC point should exist to build a bridge to triplet S. As measured in experiments, 25 % of

CS2 proceeds through the ISC to form ground state 3S [126]. The difference of electron

quantum spin number of ground state 1CS2 and 3S has prompted consideration of the energy

level of relevant reduced sulfur species. It is revealed that the ground state of species constituted

by only S and O atoms are in triplet (3SO, 3S, 3O and 3O2), except for 1SO2, while the gourd

state for species such as 1CS2, 1COS, 1CS and 1SO2 are in singlet. If we consider all oxidation

processes to have proceeded on the ground state surface, we should have the reaction process

with two ISC points as shown below:

1CS2 → 3SO → 1SO2


43

Table 2.4 summarises the branching ratio of 3S/1S from photolysis of reduced sulfur species,

as measured in experiments. Although the parent compounds are in their singlet (ground) state,

the corresponding sulfurs atom arise in both singlet and triplet states. This has indicated that

the ISC is predominant in sulfur related species, especially when it couples with O2.

Table 2.4. The branching ratio of 3S/1S from photolysis of reduced sulfur species (H2S, COS

and CS2).

Species 3S/1S Wavelength (nm) Source

H2S 0.87/0.13 193 [122]

COS 0.95/0.05 222 [123]

CS2 0.75/0.25 193 [124]

However, the identification of electronic spin states of intermediates or prod

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