Investigation of efficient two-step methanol synthesis ...

INVESTIGATION OF EFFICIENT, TWO-STEP

METHANOL SYNTHESIS PROCESSES

AT MODERATE PRESSURES AND TEMPERATURES

FAN WU, M. ENG-CHEM

ORCID: 0000-0001-5078-8711

Submitted in total fulfillment of the requirement of the degree of Doctor of Philosophy

February 2018

Department of Chemical Engineering

The University of Melbourne

Australia

ii

ABSTRACT

Methanol is one of the most common chemical commodities traded around the world

every year. It can be used as a raw material with wide applications such as construction

materials, plastics and fuel cells. Currently methanol is commercially manufactured from

syngas containing carbon monoxide (CO) and hydrogen (H2) via a one-step method at

high temperatures and pressures (250 to 300 °C and 50 to 100 bar). However, the harsh

operating conditions have negative impacts on the economy and the environment.

Compared with the conventional methanol synthesis process, a two-step method of

producing methanol in the liquid phase has attracted increasing interest as the operating

conditions are moderate (100 °C to 140 °C and 30 bar), thus reducing the operating cost.

This method consists of two steps of reactions: (1) carbonylation of alcohol to produce

ester; (2) hydrogenation reaction of the produced ester to produce methanol. The

hydrogenation reaction is the rate limiting step in the two-step methanol synthesis

process and is the focus of this work.

In this work, an extensive literature review of two-step methanol synthesis process,

including the vapour-liquid equilibrium (VLE), reactions of each step at moderate

temperatures and pressures, and commercial catalysts used for the synthesis process,

was conducted.

The solubilities of carbon monoxide (CO) and hydrogen (H2) in liquid methanol and

methyl formate were comprehensively studied in a designed apparatus from 296 K to

375 K in preparation for study the reaction properties. The solubilities of H2 in methanol

and methyl formate are lower than those of CO in methanol and methyl formate. The

solubilities of CO and H2 in methanol and methyl formate increase with increasing

iii

temperature probably due to the endothermic processes. The experimental data was

further validated using a modified Peng-Robinson Equation of State through an equation

of state approach (phi-phi method). The binary interaction parameters of the model used

for predicting the VLE were determined. The binary interaction parameters for CO in

methanol and methyl formate were constant, which were independent of the

temperature, whereas the binary interaction parameters for H2 in methanol and methyl

formate was a function of temperature.

The reaction mechanism of the hydrogenation reaction of methyl formate using a

commercial catalyst, copper chromite, was thoroughly investigated. A possible reaction

mechanism containing five elementary steps of (1) the adsorption of methyl formate on

the catalyst surface; (2) the adsorption and dissociation of hydrogen on the catalyst

surface; (3) surface reaction; (4) methanol generation via the intermediate CH2OH(s); (5)

methanol generation via the intermediate CH3O(s) was proposed. The rate constants

were estimated using MATLAB build-in functions ‘ode15’ and ‘fmincon’ through the least

squares minimisation method. The step of H2 adsorption and dissociation on the catalyst

surface was speculated as the rate controlling step due to the lowest rate constant

obtained among all forward reactions, and the corresponding activation energy was

determined as 50.15 kJ/mol, which was consistent with other published works. Through

using the evaluated rate constants, the reaction kinetics and mechanism were validated

from a group of hydrogenation reaction experiments conducted at 1.8 MPa, 2.0 MPa and

2.2 MPa with a same temperature of 110 °C.

A novel catalytic system of Cu/ZnO/ZrO2-hydrotalcite (Cu/ZnO/ZrO2-HTC) was

developed for the hydrogenation reaction of methyl formate at moderate temperatures

and pressures via a simple co-precipitation method. The ratio and role of each component

iv

were comprehensively investigated and compared using a number of characteristic

techniques including TPR, XRD, SEM, TPD-CO2, TGA, BET and XPS. An optimised catalytic

system of Cu/ZrO2(8:2)-HTC was screened, used and compared with the commercial

catalyst, copper chromite. The developed catalyst has better performance at relatively

lower temperatures, where the reaction rate of the hydrogenation reaction using the

novel catalyst in the first 100 minutes is four times faster than that of using the

commercial copper chromite, and the time to achieve equilibrium of using the novel

catalyst reduces to one third of the commercial catalyst.

Compared to the traditional methanol synthesis process, this work shows great potential

to employ a two-step methanol synthesis process via methyl formate at moderate

temperatures and pressures, thus saving energy cost and operating cost.

v

DECLARATION

This is to certify that:

i. This dissertation comprises only my original work towards to the Ph.D.

ii. Due acknowledgement has been made in the text to all other materials used

iii. The dissertation is less than 100,000 words in length, exclusive of tables, maps,

bibliographies and appendices.

Fan Wu

vi

ACKNOWLEDGEMENTS

This Ph.D. is one of the most difficult tasks I have accomplished in my life to date. It is a

special and extraordinary journey and I genuinely enjoyed my adventure over the last

four years. This thesis would have never been written and completed without the

supports and dedications from many people. I would like to take this chance to thank

those great people who helped and guided me throughout my Ph.D. life.

First and foremost, I would like to express my deepest gratitude to Prof. Paul A. Webley,

who acted as my main supervisor and project leader. He offered me an opportunity to

work with him and his team, and also opened a door for expanding my outlook on the

research in chemical engineering fields. His continuous support not only in academic

work but also in finance motivated me to keep going throughout my Ph.D. life. His

suggestions and feedback were always invaluable and made my Ph.D. journey go

smoother and with less frustration.

I also appreciate the selfless contribution from Dr. Fatin Abbas Hasan and Dr. Penny Xiao,

who acted as my co-supervisors. From their meticulous care not only in my research but

also in my daily life, I always felt the warmth and comfort like at home. Also, their rigorous

academic attitudes towards work encouraged me to be optimistic to every challenge

throughout my Ph.D. I also would like to acknowledge Dr. Luke Connel and later Prof.

David Dunstan for serving as my research committee chair, who provided me with

valuable comments on each year’s progress review.

I am also thankful to my kind and helpful colleagues in our clean energy lab, especially

David Danaci, Dr. Ranjeet Singh, Lefu Tao, Xin Fang and Thomas Moore, who helped me

set up the experiments and troubleshoot uncountable bottlenecks in my research.

vii

Without them, I may not be sitting here to write up my thesis. I am also thankful to Dr

Chaoen Li in CSIRO, who provided constructive advices in deriving the chemical reaction

mechanism. Also, I felt to be obliged to mention Prof. Sandra Kentish, Prof. Geoff Stevens,

Prof. David Shallcross, Dr. Daniel Heath, and Dr. Gabriel De Silva, who provided me with

casual tutorial positions in the department to enrich my teaching experience. I would like

to thank my research project (RP) student, Peng Shang, to come and conduct some

experiments for 12 weeks. For all my friends, especially Chen Yuan, Chloe Jack, Feng Li,

Qinghu Zhao, Yuhan Men, and Guoping Hu, their long-lasting friendship makes my work

and life more pleasant.

To Yue (Frank) Wu, my wonderful husband. He has been considerable and understanding

during my Ph.D. journey. He gave me everything he has and looked after me very well.

Despite being very busy with his Ph.D. work in the same department, he has always been

here for me, helping me through this adventure. I definitely would not have survived

during the hard time of my Ph.D. without him. I truly thank to his patience and endless

love for me.

Finally, to my beloved parents. None of this would have been possible without them. Both

have financially supported my overseas undergraduate study and living in Melbourne. I

am always number one to them and enjoy their endless love and care. I am so fortunate

to be your daughter and I am so proud of you.

viii

Contents Abstract .............................................................................................................................................................. ii

Declaration ........................................................................................................................................................ v

Acknowledgements ...................................................................................................................................... vi

List of Equations .......................................................................................................................................... xiv

List of Figures ............................................................................................................................................ xviii

List of Reactions ....................................................................................................................................... xxiii

List of Tables ................................................................................................................................................ xxv

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

1.1 Background ..................................................................................................................................... 1

1.2 Motivation, objective and outline of the thesis ................................................................. 5

Chapter 2 Literature review ................................................................................................................. 7

2.1 Commercial production of methanol .................................................................................... 7

2.1.1 Syngas production ............................................................................................................... 7

2.1.2 Methanol synthesis from syngas.................................................................................... 9

2.1.3 Separation and purification ........................................................................................... 14

2.1.4 Limitations of commercial methanol synthesis processes ................................ 16

2.2 Methanol synthesis at moderate conditions .................................................................... 18

2.2.1 Three approaches of methanol synthesis ................................................................. 18

2.2.2 Methanol synthesis via methyl formate (MS via MF) .......................................... 24

2.3 Solubility in gas-liquid phase ................................................................................................. 37

ix

2.3.1 Definition of solubility ..................................................................................................... 37

2.3.2 Fugacity of the system ..................................................................................................... 38

2.3.3 Solubility literature data review .................................................................................. 39

2.3.4 Solubility modelling .......................................................................................................... 50

2.4 Conclusions ................................................................................................................................... 57

Chapter 3 Materials and methodology ........................................................................................... 58

3.1 Materials ........................................................................................................................................ 58

3.2 Methodologies ............................................................................................................................. 59

3.2.1 Powder X-ray diffraction (XRD) ................................................................................... 59

3.2.2 Scanning electron microscopy (SEM) ........................................................................ 59

3.2.3 Energy Dispersive X-ray spectroscopy (EDX) ........................................................ 59

3.2.4 N2 adsorption-desorption isotherms ......................................................................... 59

3.2.5 Temperature-programmed reduction (TPR) .......................................................... 60

3.2.6 Specific copper surface area via N2O titration ........................................................ 62

3.2.7 Thermal gravimetric analysis (TGA) .......................................................................... 63

3.2.8 X-ray photoelectron spectroscopy (XPS).................................................................. 64

3.2.9 Auger Electron Spectroscopy (AES) ........................................................................... 64

3.2.10 Products analysis method - Gas chromatography (GC) .................................. 64

Chapter 4 Solubility study ................................................................................................................... 67

4.1 Objective ........................................................................................................................................ 67

4.2 Experimental apparatus and procedures.......................................................................... 67

x

4.2.1 Apparatus .............................................................................................................................. 67

4.2.2 Procedure .............................................................................................................................. 69

4.3 Theory ............................................................................................................................................. 72

4.3.1 Evaluation of experimental results ............................................................................. 72

4.3.2 Modelling .............................................................................................................................. 74

4.3.3 Uncertainty calculation ................................................................................................... 79

4.3.4 Henry’s law constant and its confidence intervals ............................................... 80

4.3.5 Thermodynamic property determination ................................................................ 81

4.4 Data analysis ................................................................................................................................. 82

4.4.1 Validation of the experimental apparatus................................................................ 82

4.4.2 Experimental results ........................................................................................................ 85

4.4.3 Modelling results ............................................................................................................... 93

4.5 Conclusions ................................................................................................................................... 98

Chapter 5 Hydrogenation reaction kinetics mechanism ......................................................... 99

5.1 Introduction .................................................................................................................................. 99

5.2 Experimental apparatus and procedures....................................................................... 102

5.2.1 Apparatus ........................................................................................................................... 102

5.2.2 Procedure ........................................................................................................................... 103

5.3 Carbonylation reaction study ............................................................................................. 106

5.4 Hydrogenation reaction catalysts preparation and characterisation ................. 107

5.4.1 Catalysts preparation .................................................................................................... 107

xi

5.4.2 The structure and phase compositions .................................................................. 108

5.4.3 Thermal behaviour and stability .............................................................................. 110

5.4.4 The morphology and size............................................................................................. 111

5.4.5 The surface area and specific copper surface area ............................................ 112

5.4.6 Summary ............................................................................................................................ 113

5.5 Study of hydrogenation reaction kinetics and explore the reaction mechanism ...

........................................................................................................................................................ 113

5.5.1 Effects of agitation speed on reaction rate ........................................................... 113

5.5.2 Effects of catalysts loadings on reaction rate ....................................................... 115

5.5.3 Effects of temperature on reaction conversion and selectivity .................... 117

5.5.4 Reaction kinetics model and parameter estimation ......................................... 122

5.5.5 Mechanism validation ................................................................................................... 132

5.6 Conclusions ................................................................................................................................ 138

Chapter 6 Development of novel hydrogenation catalysts ................................................. 140

6.1 Introduction ............................................................................................................................... 140

6.2 Design of a novel catalyst ..................................................................................................... 141

6.3 Cu/Zn/Zr-HTC catalyst .......................................................................................................... 143

6.3.1 Catalyst preparation ...................................................................................................... 143

6.3.2 Catalyst reducibility ....................................................................................................... 145

6.3.3 Catalyst crystalline structure ..................................................................................... 147

6.3.4 Thermal stability of the catalyst ............................................................................... 150

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6.3.5 Morphology and dispersion of the catalyst .......................................................... 153

6.3.6 Surface areas and specific copper surface area .................................................. 157

6.3.7 The surface basicity of CuZnZr-HTC catalyst ....................................................... 158

6.3.8 Chemical states of elements in the catalyst .......................................................... 160

6.4 Conclusions ................................................................................................................................ 171

Chapter 7 Determination of the catalytic performance of the novel hydrogenation

catalysts ................................................................................................................................................ 172

7.1 Introduction ............................................................................................................................... 172

7.2 Experimental apparatus and procedures....................................................................... 173

7.3 Results and Discussion .......................................................................................................... 173

7.3.1 Roles of the components in the Cu/ZnO/ZrO2-HTC catalytic system......... 173

7.3.2 By products formation .................................................................................................. 176

7.3.3 The catalytic effect of the ratio of Cu/ZnO on the reaction ............................ 177

7.4 Comparison of Cu/ZrO2-HTC catalysts with copper chromite............................... 182

7.4.1 The characteristics of catalysts ................................................................................. 182

7.4.2 Catalytic performance ................................................................................................... 183

7.5 Conclusions ................................................................................................................................ 186

Chapter 8 Conclusions and recommendations ......................................................................... 188

Chapter 9 References ......................................................................................................................... 190

Chapter 10 Appendices .................................................................................................................... 214

Appendix A Mass Spectrometry (MS) calibration of hydrogen using 5.4% H2/Ar ..... 214

xiii

Appendix B. Mass Spectrometry (MS) calibration of CO2 using 4.99% CO2/He .......... 215

Appendix C. Multi-level calibration of liquid samples in GC ................................................ 216

Appendix D. The grade of the certified standard-spec gas from ScottTM for retention

time determination .............................................................................................................................. 217

xiv

LIST OF EQUATIONS

Equation 2-1. Formula for S value determination ................................................................................................. 10

Equation 2-2. The overall reaction rate expression [86] ....................................................................................... 26

Equation 2-3. The overall reaction rate expression [87] ....................................................................................... 27

Equation 2-4. Carbonylation reaction forward reaction rate [88] ........................................................................ 27

Equation 2-5. Carbonylation reaction forward reaction rate with pyridine as a promoter [88]........................... 27

Equation 2-6. Reaction kinetics expression for the hydrogenation reaction [65] ................................................. 30

Equation 2-7. Reaction kinetics expression for the hydrogenation reaction in the presence of CO [65] .............. 30

Equation 2-8. Reaction kinetics expression for the hydrogenation reaction [86] ................................................. 30

Equation 2-9. Predicted concurrent methanol production [89] ............................................................................ 33

Equation 2-10.Thermodynamic equilibrium of the fugacity in gas phase and liquid phase ................................. 37

Equation 2-11. The equation of the fugacity coefficient ....................................................................................... 38

Equation 2-12. The equation of activity coefficient .............................................................................................. 39

Equation 2-13. Krichevsky-Kasarnovsky equation ................................................................................................ 43

Equation 2-14. Generalised equation of Peng Robinson EoS ................................................................................ 51

Equation 2-15. The parameter b determination of binary system........................................................................ 51

Equation 2-16. The parameter a determination of binary system........................................................................ 51

Equation 2-17. b of the pure component .............................................................................................................. 51

Equation 2-18. a of the pure component .............................................................................................................. 51

Equation 2-19. Alpha function by Soave ............................................................................................................... 52

Equation 2-20. Generalised function m of acentric functor .................................................................................. 52

Equation 2-21. Alpha function by Twu .................................................................................................................. 52

Equation 2-22. Generalised alpha function by Twu .............................................................................................. 53

Equation 2-23. Alpha function by Boston-Mathias at subcritical condition ......................................................... 53

Equation 2-24. Alpha function by Boston-Mathias at supercritical condition ...................................................... 53

Equation 2-25. Binary interaction parameter (BIP) .............................................................................................. 54

Equation 2-26. Absolute average deviation relatives (AADR) ............................................................................... 55

xv

Equation 2-27. Fugacity equilibrium ..................................................................................................................... 55

Equation 2-28. Vapour phase fugacity of component .......................................................................................... 56

Equation 2-29. Liquid phase fugacity .................................................................................................................... 56

Equation 2-30. Liquid phase fugacity in terms of Henry’s law constant ............................................................... 56

Equation 2-31. The relationship between gas phase fugacity and liquid phase activity coefficient..................... 56

Equation 2-32. Pseudo Henry’s law constant ....................................................................................................... 56

Equation 2-33. The formula of pseudo Henry’s law constant ............................................................................... 56

Equation 3-1. Determination of amount of desorbed CO2 .................................................................................... 62

Equation 3-2. Determination of copper dispersion ............................................................................................... 63

Equation 4-1. The total amount of solute in the tank 4 ........................................................................................ 72

Equation 4-2. Generalised equation of Peng-Robinson EoS.................................................................................. 72

Equation 4-3. Parameter b in PR EoS .................................................................................................................... 72

Equation 4-4. Parameter a in PR EoS .................................................................................................................... 72

Equation 4-5. alpha function ................................................................................................................................ 73

Equation 4-6. Determination of mi when acentric factor less than 0.49 .............................................................. 73

Equation 4-7. Determination of mi when acentric factor above 0.49 .................................................................. 73

Equation 4-8. Total moles of solvents in the equilibrium cell ............................................................................... 73

Equation 4-9. The density of the solvent .............................................................................................................. 73

Equation 4-10. The total pressure in the absorption tank 4 ................................................................................. 74

Equation 4-11. The Antoine Equation ................................................................................................................... 74

Equation 4-12. The total amount of the solvent in the tank 4 .............................................................................. 74

Equation 4-13. The total amount of the solute in the tank 4 ................................................................................ 74

Equation 4-14. The amount of solvent in the vapour phase ................................................................................. 74

Equation 4-15. The mole fraction of solute in solvent .......................................................................................... 74

Equation 4-16. Determination of b parameter in PR-EoS ..................................................................................... 75

Equation 4-17. Determination of a parameter in PR-EoS ..................................................................................... 75

Equation 4-18. Determination of bi parameter in PR-EoS .................................................................................... 75

Equation 4-19. Determination of ai parameter in PR-EoS .................................................................................... 75

xvi

Equation 4-20. Fugacity equilibrium ..................................................................................................................... 75

Equation 4-21. Wilson’s equation ......................................................................................................................... 77

Equation 4-22. The expression of the mole fraction of components in the liquid phase ...................................... 77

Equation 4-23. The expression of the mole fraction of components in the gas phase ......................................... 77

Equation 4-24. Determination of the mole fraction of components in terms of equilibrium ratio K .................... 77

Equation 4-25. Fugacity coefficient of components in the liquid phase ............................................................... 78

Equation 4-26. Determination of the mixture parameter Ѱ in the liquid phase .................................................. 78

Equation 4-27. Determination of the mixture parameter 𝑎𝛼 in the liquid phase ................................................. 78

Equation 4-28. Determination of the mixture parameter b in the liquid phase ................................................... 78

Equation 4-29. Fugacity coefficient of components in the gas phase .................................................................. 78

Equation 4-30. Determination of the mixture parameter Ѱ in the gas phase ...................................................... 78

Equation 4-31. Determination of the mixture parameter 𝑎𝛼 in the gas phase .................................................... 78

Equation 4-32. Determination of the mixture parameter b in the gas phase ....................................................... 79

Equation 4-33. Evaluation of a new equilibrium ratio K ....................................................................................... 79

Equation 4-34. Convergence constrains ............................................................................................................... 79

Equation 4-35. Uncertainty of u(x)/x .................................................................................................................... 80

Equation 4-36. Uncertainty of u(n)/n in the gas phase......................................................................................... 80

Equation 4-37. Uncertainty of u(n)/n in the liquid phase ..................................................................................... 80

Equation 4-38. The Henry’s constant expression .................................................................................................. 80

Equation 4-39. The fugacity of the solute in the liquid phase ............................................................................... 80

Equation 4-40. The expression of the Henry’s constant ........................................................................................ 81

Equation 4-41. Dissolution enthalpy of gas-liquid solubility ................................................................................. 81

Equation 4-42. Dissolution entropy of gas-liquid solubility .................................................................................. 81

Equation 4-43. Dissolution Gibbs free energy of gas-liquid solubility ................................................................... 81

Equation 4-44. The binary interaction parameter ................................................................................................ 95

Equation 4-45. Definition of AARD ........................................................................................................................ 95

Equation 5-1. Reaction rate expression of methyl formate ................................................................................ 126

Equation 5-2. Reaction rate expression of H2 ..................................................................................................... 126

xvii

Equation 5-3.Reaction rate expression of methanol .......................................................................................... 126

Equation 5-4. Reaction rate expression of HCOOCH3(s2) .................................................................................... 126

Equation 5-5. Reaction rate expression of H(s) ................................................................................................... 126

Equation 5-6. Reaction rate expression of CH2OH(s) .......................................................................................... 127

Equation 5-7. Reaction rate expression of CH3O(s) ............................................................................................. 127

Equation 5-8. Reaction rate expression of catalytic site s .................................................................................. 127

Equation 5-9 BDF evaluation using higher orders of Taylor polynomial ............................................................. 127

Equation 5-10. Least-squares regression function .............................................................................................. 127

Equation 5-11. The absolute average relative residual ...................................................................................... 128

Equation 5-12. Central differences ..................................................................................................................... 129

Equation 5-13 Precision matrix P ........................................................................................................................ 129

Equation 5-14 Degrees of freedom ..................................................................................................................... 129

Equation 5-15 Residual variance ........................................................................................................................ 130

Equation 5-16. Arrhenius equation ..................................................................................................................... 130

Equation 6-1. Auger parameter (αCu) .................................................................................................................. 163

Equation 7-1. Pseudo-space time yield (STYPS) ................................................................................................... 178

xviii

LIST OF FIGURES

Figure 1-1. Methanol as the basic chemical and energy feedstock [10] ................................................................. 2

Figure 1-2. Global methanol demand by end-use 2015 [14] .................................................................................. 3

Figure 1-3. Global methanol supply and demand [13] ........................................................................................... 4

Figure 1-4. World methanol demand by regions [28] ............................................................................................. 4

Figure 2-1. Global methanol production technologies distribution ...................................................................... 11

Figure 2-2. Thermodynamic equilibrium curves for the conversion of syngas to methanol. Syngas with

(H2/CO=2) is used .................................................................................................................................................. 17

Figure 2-3. Equilibrium conversion of CO in methanol synthesis at various temperatures and pressures. Syngas

with (H2/CO=2) is used .......................................................................................................................................... 17

Figure 2-4. A proposed mechanism of methanol synthesis using low grade syngas based on the DRIFT study ... 21

Figure 2-5. Thermodynamic analysis of methanol reaction condition when H2/CO = 2 [89] ................................ 33

Figure 2-6. Literature data of H2-methanol binary system at room temperature at low pressures. The x-axis

stands for mole fraction of H2. The y-axis stands for total pressure. (+) Choudhary et al. at 293 K [100]; (*)

Wainwright et al. at 291 K [101]; (∆) Descamps et al. at 291.2 K [108]; (○) Liu et al. at 296.25 K [103] ............ 43

Figure 2-7. Literature data of H2-methanol binary system at room temperature at high pressures. The x-axis

stands for mole fraction of H2. The y-axis stands for total pressure. (squares) Brunner et al. at 298.15 K [106];

(solid circles) Bezanehtak et al. at 298.15 K [104] ................................................................................................ 44

Figure 2-8. Literature data of H2-methanol binary system at various temperature at low pressure. (X) Liu et al.

at 373.95 K [103]; (solid square) Liu et al. at 363.55 K [103];(Φ) Descamps et al. at 308.2 K [108] .................... 45

Figure 2-9. Literature data of H2-methanol binary system at various temperature at high pressure [106] ......... 45

Figure 2-10. Solubility data of H2-methyl formate binary system. The x-axis stands for the mole fraction of H2.

The y-axis stands for the total pressure. (a) by Liu et al. [103]; (b) by Wainwright et al. [101]........................... 47

Figure 2-11. Solubility data of CO-methanol binary system at 323 K. The x-axis stands for the mole fraction of

H2. The y-axis stands for the total pressure. (x) Liu et al. [111]; (○)Tonner et al.[110]; (∆) Brunner et al. [106] 49

Figure 2-12. Solubility data of CO-methanol binary system at various temperatures [106]. The x-axis stands for

the mole fraction of H2. The y-axis stands for the total pressure. ........................................................................ 50

xix

Figure 3-1. Schematic diagram of the valves and detectors in the Agilent GC 7890B .......................................... 66

Figure 4-1. Schematic diagram of solubility apparatus: 1. Gas cylinders (He, CO2, H2 and CO); 2. Mass flow

controller, 3. Storage tank; 4. Absorption tank; 5. Heating tape/Cooling bath; 6. Magnetic stirrer; 7. Vacuum

pump; 8. Vent system; BV-1 to BV-4: Ball valves; NV-1 and NV-2: Needle valves ................................................ 68

Figure 4-2. The flow chart of solubility experiments ............................................................................................. 71

Figure 4-3. The flow chart of phi-phi approach to determine the VLE data ......................................................... 77

Figure 4-4. The comparison of the experimental results with the literature data ................................................ 83

Figure 4-5. The comparison of the experimental results and literature data for CO-CH3OH system at 298.1 K... 84

Figure 4-6. The composition of the experimental results and literature data for CO-CH3OH system at 322.7 K .. 85

Figure 4-7. Isothermal phase equilibrium of CO in methyl formate...................................................................... 87

Figure 4-8. Isothermal phase equilibrium of CO in methyl formate...................................................................... 89

Figure 4-9. Isothermal phase equilibrium of H2 in methanol ................................................................................ 90

Figure 4-10. Isothermal phase equilibrium of H2 in methyl formate .................................................................... 91

Figure 4-11. Modelling validation results of CO solubility in methanol ................................................................ 96

Figure 4-12. Modelling validation results of CO solubility in methyl foramte ...................................................... 96

Figure 4-13. Modelling validation results of H2 solubility in methanol ................................................................. 97

Figure 4-14. Modelling validation results of H2 solubility in methyl formate ....................................................... 97

Figure 5-1. Schematic diagram of reaction apparatus: 1. Gas cylinders (He, CO2, H2 and CO); 2. Mass flow

controller, 3. Storage tank; 4. Reactor; 5. Heating tape/Cooling bath; 6. Magnetic stirrer; 7. Vacuum pump; 8.

Vent system; 9. Gas sampling tank; 10. GC; BV-1 to BV-6: Ball valves; NV-1 to NV-3: Needle valves ................ 102

Figure 5-2. The pressure profiles of carbonylation reaction. Operating conditions: Ptotal=2.3 MPa, agitation

speed = 800 rpm, catalyst loadings = 0.4 mol/L ................................................................................................. 107

Figure 5-3. The TPR profile of the copper chromite catalysts ............................................................................. 108

Figure 5-4. XRD patterns of copper chromite catalyst ........................................................................................ 110

Figure 5-5. The profile of thermal gravity analysis and the corresponding DTG ................................................ 111

Figure 5-6. SEM images of the copper chromite sample. (a) and (b) are copper chromite; (c) and (d) are reduced

copper chromite .................................................................................................................................................. 112

xx

Figure 5-7. Effect of rotation speeds on the conversion of methanol. Operating conditions: Ptotal = 3.2 MPa, T =

384 K, Catalyst loading = 16 g/L ......................................................................................................................... 114

Figure 5-8. Effect of the catalyst loadings on the conversion rate of methanol. Operating conditions: Ptotal= 3.2

MPa, T = 384K. Rotation speed = 800 rpm. ........................................................................................................ 116

Figure 5-9. Effect of temperature on the reaction rate. Operating conditions: Ptotal = 3.2 MPa, T = 346 K, Catalyst

loading = 16 g/L. Rotation speed = 800 rpm. ...................................................................................................... 118

Figure 5-10. Effect of temperature on the reaction rate. Operating conditions: Ptotal = 3.2 MPa, T = 370 K,

Catalyst loading = 16 g/L. Rotation speed = 800 rpm. ........................................................................................ 119

Figure 5-11. Effect of temperature on the reaction rate. Operating conditions: Ptotal = 3.2 MPa, T = 384 K,

Catalyst loading = 16 g/L. Rotation speed = 800 rpm. ........................................................................................ 120

Figure 5-12. Comparison of experimental and simulation results at T = 346 K. Operating conditions: Ptotal = 3.2

MPa, Catalyst loading = 16 g/L. Rotation speed = 800 rpm. .............................................................................. 131





Figure 5-15. Validation of the modelling parameters on various pressure. Catalyst loading = 16 g/L, T = 384K.

Operating conditions: PH2 = 1.8 MPa, T = 384 K, Catalyst loading = 16 g/L. Rotation speed = 800 rpm ............. 133

Figure 5-16. Validation of the modelling parameters on various pressure. Operating conditions: PH2 = 2.0 MPa, T

= 384 K, Catalyst loading = 16 g/L. Rotation speed = 800 rpm ........................................................................... 133

Figure 5-17. Validation of the modelling parameter on various pressure. Operating conditions: PH2 = 2.2 MPa, T

= 384 K, Catalyst loading = 16 g/L. Rotation speed = 800 rpm ........................................................................... 134

Figure 5-18. Schematic description of the proposed mechanism. Step 1 to Step 3 ............................................ 136

Figure 5-19. Schematic description of the proposed mechanism. Step 4 and Step 5 ......................................... 137

Figure 6-1. TPR profile of calcinated CuZnZr-HTC catalysts with different Cu/Zn ratio. (a) Cu8 (b) Cu6 (c) Cu4 (d)

Cu2 ...................................................................................................................................................................... 146

Figure 6-2. XRD patterns of dried CuZnZr-HTC catalyst ...................................................................................... 148

Figure 6-3. XRD patterns of calcinated CuZnZr-HTC catalyst .............................................................................. 149

xxi

Figure 6-4. XRD patterns of reduced CuZnZr-HTC catalyst ................................................................................. 150

Figure 6-5. Thermogravimetry profiles of dried CuZnZr catalysts. (a) activated HTC; (b) dCu0; (c) dCu2; (d) dCu4;

(e) dCu6; (f) dCu8 ................................................................................................................................................ 153

Figure 6-6. SEM images and mapping of the rCu0 sample ................................................................................. 154




Figure 6-10. SEM images and mapping of the rCu8 sample ............................................................................... 156

Figure 6-11. CO2-TPD pattern of the reduced CuZnZr-HTC catalysts. (a) Cu8; (b) Cu6; (c) Cu4; (d) Cu2 ............. 160

Figure 6-12. Cu 2p core level X-ray photoelectron spectra of CnZuZr-HTC series samples. (A) Cu2; (B) Cu4; (C)

Cu6; (D) Cu8 (i) represents calcinated state, (ii) represents reduced state. ........................................................ 166

Figure 6-13. Zn 2p core level X-ray photoelectron spectra of CnZuZr-HTC series samples. (A) Cu2; (B) Cu4; (C)


Figure 6-14. Zr 3d core level X-ray photoelectron spectra of CnZuZr-HTC series samples. (A) Cu2; (B) Cu4; (C)


Figure 6-15. O 1s core level X-ray photoelectron spectra of CnZuZr-HTC series samples. (A) Cu2; (B) Cu4; (C) Cu6;

(D) Cu8 (i) represents calcinated state, (ii) represents reduced state. ................................................................ 169

Figure 6-16. X-ray induced Auger electron spectra of catalysts. (a) cCu2; (b) rCu2; (c) cCu4; (d) rCu4; (e) cCu6; (f)

rCu6; (g) cCu8; (h) rCu8 ....................................................................................................................................... 170

Figure 7-1. The pressure profile of the catalysts. The total pressure of the system, including the partial pressure

of the solvent and the pressure of the gas.......................................................................................................... 174

Figure 7-2. The appearance of the Cu/ZrO2 catalysts after reaction .................................................................. 175

Figure 7-3. The amount of methanol produced over time .................................................................................. 179

Figure 7-4. The amount of hydrogen in the reactor over time ........................................................................... 179

Figure 7-5. The relationship between the space time yield and copper surface area ......................................... 181

Figure 7-6. Amount of methanol and H2 in the reactor with two catalysts system. Operating conditions: Ptotal =

3.2 MPa, T = 384 K, stirrer speed: 800 rpm, catalyst loading: 16 g/L ................................................................. 184

xxii

Figure 7-7. The total pressure profiles from two catalysts. Operating conditions: Ptotal = 3.2 MPa, T = 370 K,

stirrer speed: 800 rpm, catalyst loading: 16 g/L ................................................................................................. 185

Figure 7-8. Amount of methanol and H2 in the reactor with two catalysts system. Operating conditions: Ptotal =

3.2 MPa, T = 370 K, stirrer speed: 800 rpm, catalyst loading: 16 g/L ................................................................. 185

xxiii

LIST OF REACTIONS

Reaction 2-1. Methane steam reforming ............................................................................................................... 7

Reaction 2-2. Water gas shift reaction (WGS) ........................................................................................................ 8

Reaction 2-3. Methane partial oxidation ................................................................................................................ 8

Reaction 2-4. Oxidation of carbon monoxide ......................................................................................................... 8

Reaction 2-5. Oxidation of hydrogen ...................................................................................................................... 8

Reaction 2-6. Oxidation of carbon .......................................................................................................................... 8

Reaction 2-7. Reaction with carbon and water ...................................................................................................... 8

Reaction 2-8. Water gas shift reaction ................................................................................................................... 8

Reaction 2-9. Formation of carbon monoxide via carbon dioxide and carbon ....................................................... 8

Reaction 2-10. Methanol synthesis I ....................................................................................................................... 9

Reaction 2-11. Methanol synthesis II ...................................................................................................................... 9

Reaction 2-12. Reverse water gas shift reaction .................................................................................................. 10

Reaction 2-13. Carbonylation reaction of alcohol ................................................................................................ 21

Reaction 2-14. Hydrogenation reaction of ester ................................................................................................... 21

Reaction 2-15. Mechanism step 1 ......................................................................................................................... 25

Reaction 2-16. Mechanism step 2 ......................................................................................................................... 25

Reaction 2-17. Formation of sodium formate [90] ............................................................................................... 28

Reaction 2-18. Formation of sodium formate [90] ............................................................................................... 29

Reaction 2-19. Catalyst deactivation by CO2 ........................................................................................................ 29

Reaction 2-20. De-carbonylation reaction ............................................................................................................ 31

Reaction 2-21. Carbonylation reaction ................................................................................................................. 39

Reaction 2-22. Hydrogenation reaction ................................................................................................................ 39

Reaction 3-1. Reduction of copper (II) oxide to metallic copper ........................................................................... 62

Reaction 3-2. Oxidation of copper by nitrous oxide to copper (I) oxide ................................................................ 62

Reaction 3-3. Reduction of copper (I) oxide to metallic copper ............................................................................ 63

Reaction 5-1. Hydrogenation of methyl formate ................................................................................................ 100

xxiv

Reaction 5-2. Decomposition of copper barium ammonium chromite ............................................................... 101

Reaction 5-3. Decomposition of copper ammonium chromite ........................................................................... 101

Reaction 5-4. Decarbonylation reaction ............................................................................................................. 122

Reaction 5-5. Dehydration of methanol ............................................................................................................. 122

Reaction 5-6. Adsorption of methyl formate on the catalyst active sites ........................................................... 123

Reaction 5-7. Adsorption of H2 on the catalyst active sites ................................................................................ 123

Reaction 5-8. Formation of intermediates .......................................................................................................... 123

Reaction 5-9. Production of methanol I .............................................................................................................. 123

Reaction 5-10. Production of methanol II ........................................................................................................... 123

Reaction 7-1. Decarbonylation reaction of methyl formate ............................................................................... 176

Reaction 7-2. Dehydration of methanol ............................................................................................................. 176

xxv

LIST OF TABLES

Table 2-1. The current methanol synthesis suppliers with their methanol convertors/reactors .......................... 12

Table 2-2. Selection of methanol distillation trains .............................................................................................. 15

Table 2-3. The change in free energy (∆𝑮) of methanol synthesis at various temperatures ................................ 17

Table 2-4. Comparison of three types of non-conventional methanol synthesis processes with conventional

methanol synthesis ............................................................................................................................................... 23

Table 2-5. Published literature for H2-CH3OH binary system ................................................................................ 40

Table 2-6. The generalized Twu alpha function parameters for subcritical and supercritical conditions ............. 53

Table 2-7. Coefficients of two parameter and three parameter equations in H2-methanol system ..................... 55

Table 3-1. Information of chemicals used in the study ......................................................................................... 58

Table 3-2. Information of gas cylinders used in the study .................................................................................... 58

Table 4-1. Physical properties of pure components .............................................................................................. 73

Table 4-2. The parameters for the solvent density determination ....................................................................... 73

Table 4-3. The parameters for the solvent saturated pressure ............................................................................ 74

Table 4-4. Partial pressure (PCO2), liquid phase mole fraction (xi), and uncertainties (δ) of CO2 in methanol from

298.15 K to 373.15 K ............................................................................................................................................. 82

Table 4-5. Partial pressure (PCO), liquid phase mole fraction (xi), Henry’s law constant (H) and uncertainties (δ)

of CO in methanol from 298.15 K to 373.15 K ...................................................................................................... 87

Table 4-6. Partial pressure (PCO), liquid phase mole fraction (xi), Henry’s law constant (H) and uncertainties (δ)

of CO in methyl formate from 298.15 K to 373.15 K ............................................................................................. 88

Table 4-7. Partial pressure (PH2), liquid phase mole fraction (xi), Henry’s law constant (H) and uncertainties (δ)

of H2 in methanol from 298.15 K to 373.15 K ....................................................................................................... 90

Table 4-8. Partial pressure (PH2), liquid phase mole fraction (xi), Henry’s law constant (H) and uncertainties (δ)

of H2 in methyl formate from 298.15K to 373.15 K .............................................................................................. 91

Table 4-9. The thermodynamic properties of the systems .................................................................................... 93

Table 4-10. The regressed binary parameter using PR EoS for different systems ................................................ 95

Table 4-11. Empirical coefficients of binary interaction parameters 𝑘𝑖𝑗 .............................................................. 98

xxvi

Table 5-1. Experimental conditions for hydrogenation reactions ....................................................................... 105

Table 5-2. The carbonylation reaction performance .......................................................................................... 107

Table 5-3. The reducibility of the copper chromite ............................................................................................. 108

Table 5-4. Surface properties of the copper chromite ........................................................................................ 113

Table 5-5. Amount of methanol produced under different agitation speeds .................................................... 115

Table 5-6. Amount of hydrogen at various catalyst loadings ............................................................................ 117

Table 5-7. Experimental values of reactants and products at different temperature. ....................................... 121

Table 5-8. The absolute average relative residual (AARD,%) for each system ................................................... 128

Table 5-9. Regressed kinetics parameters .......................................................................................................... 130

Table 5-10. The absolute average relative residual (AARD%) for each system .................................................. 134

Table 5-11. Experimental values of reactants and products at different Hydrogen pressure. ........................... 135

Table 6-1. Metal composition of prepared catalysts .......................................................................................... 144

Table 6-2. Centre of reduction peaks and corresponding concentrations to the TPR pattern over CuZnZr-HTC

catalysts with different Cu/Zn ratio .................................................................................................................... 146

Table 6-3. Total mass loss of the dried catalysts ................................................................................................ 152

Table 6-4. The relative surface concentration of metal (atomic %) on the CuZnZr-HTC catalysts. The values in the

parentheses are the nominal concentration normalized to the total metal content of the prepared samples .. 157

Table 6-5. Physicochemical properties of the calcinated samples with different Cu/Zn ratio. ........................... 157

Table 6-6. The amount and distribution of basic sites of the reduced CuZnZr-HTC catalysts ............................. 159

Table 6-7. XPS parameters of Cu core level in CuZnZr-HTC catalysts .................................................................. 162

Table 6-8. XPS parameters of Zn core level in CuZnZr-HTC catalysts .................................................................. 162

Table 6-9. XPS parameters of Zr core level in CuZnZr-HTC catalysts................................................................... 162

Table 6-10. XPS parameters of calcined CuZnZr-HTC samples ............................................................................ 164

Table 6-11. XPS parameters of reduced CuZnZr-HTC samples ............................................................................ 164

Table 7-1. Experiment Operating conditions ...................................................................................................... 173

Table 7-2. Metal compositions of prepared catalysts ......................................................................................... 173

Table 7-3. Catalytic performance for hydrogenation of methyl formate ........................................................... 178

Table 7-4. Physicochemical properties of copper chromite and Cu8 catalysts ................................................... 182

xxvii

Table 7-5. Surface composition of the catalyst ................................................................................................... 183

Table 7-6. Space time yield of the catalysts at 384 K .......................................................................................... 184

Table 7-7. Space time yield of the catalysts at 370 K .......................................................................................... 185

Introduction

1

CHAPTER 1 INTRODUCTION

1.1 BACKGROUND

Methanol is an important feedstock used as a fuel and solvent in various industries. It was

first obtained by Robert Boyle in 1661 as a by-product in the production of charcoal via

wood distillation, thus it is so called wood alcohol [1]. The elemental composition of

methanol remained unidentified until 1834 when Dumas and Peligot introduced the

terminology methyl alcohol [2]. Until 1923, methanol production rate was very limited,

approximately 10 - 20 L per ton of wood treated for charcoal manufacturing. Initially,

methanol was produced for the purpose of lighting, cooking and heating; however, it was

rapidly substituted by more economical fuels, such as kerosene [3].

In 1905, Sabatier suggested the first synthetic pathway for methanol production from CO

and H2 [4]. Based on those findings, in early 1920s, Mittasch et al. at BASF (Badische

Anilin and Soda Fabrik) synthesised organic oxygenates, including methanol, from syngas

that was supplied from coal gasification in the course of development of ammonia

synthesis [5]. In 1923, BASF (in Leuna, Germany) started to commercialise syngas-to-

methanol process with utilising sulphur resistant zinc chromite (ZnO-Cr2O3) catalysts

and operating conditions of 593 – 723 K and 25 - 35 MPa [6].

Due to the harsh operating conditions of methanol production operated by BASF, new

technology and catalysts were required to make the process feasible and economical.

With the invention of steam reforming of methane, sulphur free syngas was produced. In

1966, ICI (Imperial Chemical Industries, Great Britain) successfully produced methanol

at lower pressure and temperature (10 MPa and 573 K) by using a quench reactor loaded

with high activity Cu/ZnO catalyst [7]. Meanwhile, Lurgi Gesellschaft fur warme and

Introduction

2

Chemoteknik from Germany developed a process with lower operating temperature (503

-523 K) and pressure (4 – 5 MPa) via using a tubular reactor cooled with boiling water

[6].

Methanol can be made from wide range of feed stocks and has therefore become the most

commonly produced chemical at industrial scale. It is a primary liquid petrochemical

which can be utilised as a fuel and solvent. It is also a key industrially-derivable feedstock

that can be converted into various value-added products [3], [8], [9]. The feedstock of

methanol synthesis and its downstream products can be found in Figure 1-1.

Figure 1-1. Methanol as the basic chemical and energy feedstock [10]

Figure 1-2 shows an overview of methanol demand by end-use in 2015. Formaldehyde is

the largest single methanol derivative and is a key component for construction products

as well as car manufacturing [11]. Other typical methanol derivatives include acetic acids

and MTBE which account for a certain amount of end-use of methanol. However, recently,

Introduction

3

newer products such as light-olefins using MTO (methanol to olefins) and dimethyl ether

(DME), are changing the methanol applications palette [11]–[13].

Figure 1-2. Global methanol demand by end-use 2015 [14]

Over 90 methanol plants are distributed globally with an annual capacity of 120 million

metric tons [15]–[27]. About 0.2 million tons of methanol are consumed daily as chemical

feedstock and/or as fuels [28]. In 2010, the global demand for methanol reached 49

million metric tons (MMT), and is expected to exceed 95 MMT by 2020 based on the

global methanol supply-demand chart shown on Figure 1-3 and Figure 1-4. In addition,

the current dominant player in the methanol global market is China, which is driven by

the significant economic growth in the last two decades. China dominates about 54 % of

world demand, due to substantial demand for olefins which are derived from methanol

via MTO technology [11], [12], [14]. In 2020, northeast Asia led by China will dominate

70 % of global market demand, followed by North America at just 9 % and Western

Europe at 8 % [13].

7%

4%

18%2%3%

8%

3%

9%

2%

8%

9%

27%

Formaldehyde

Acetic Acid

MTBE/TAME

Methyl Methacrylate

Gasoline/Fuel

Biodiesel

Dimethyl ether

Methylamines

Chloromethanes

MTO

Solvents

Others

Introduction

4

Figure 1-3. Global methanol supply and demand [13]

Figure 1-4. World methanol demand by regions [28]

Introduction

5

1.2 MOTIVATION, OBJECTIVE AND OUTLINE OF THE THESIS

Methanol is commercially produced from fossil fuel-based syngas that contains CO and

H2 at 250 to 300 °C and 50 to 100 bar. The high pressure and temperature requirement

for syngas to methanol process has a negative impact on both economy and environment.

Over the past decades, researchers have been looking thoroughly for new catalysts which

can convert syngas to methanol at low temperature and low pressure and designing new

processes with moderate operating conditions. A methanol synthesis via methyl formate

process (MS via MF) was selected (as one of many options) for study in this thesis as it

has certain advantages, such as that it can be potentially industrialised at moderate

pressures and temperatures (100 - 140 °C and 20 - 40 bar). The process consists of two

reactions, a carbonylation reaction and a hydrogenation reaction. Based on the literature

study, the hydrogenation reaction is the rate limiting step, and hence, it becomes the

research focus of this study.

This thesis is presented in eight chapters. In Chapter 1, a brief introduction of methanol

and methanol synthesis background is introduced. Chapter 2 reviews and summarises

the published literature of the commercial methanol synthesis as well as the MS via MF

process. A short review of the vapour liquid equilibrium of the H2/CO system in methanol

and methyl formate system is also included in Chapter 2. Chapter 3 introduces the

materials and methodologies used in the current research. Chapter 4 reports on the

vapour liquid equilibrium of the reactants at the operating conditions of the reaction by

using a bench-scale custom-built apparatus. The Peng Robinson Equation of State (PR-

EoS) with a binary interaction parameter kij was fit to the data to ensure that the

equation of state (EoS) describes the experimental results correctly. In Chapter 5, the

hydrogenation reaction is performed at different temperatures and pressures and a

Introduction

6

reaction mechanism is proposed and validated. Chapter 6 and 7 present a study of a novel

catalyst synthesis using copper oxide, zinc oxide and zirconium oxide deposited on

hydrotalcite-like compounds, detailing the characterisation of the catalysts as well as the

hydrogenation reaction performance using the novel catalyst. Chapter 8 summarises the

PhD work and the key findings, together with some recommendations for future

development.

Literature review

7

CHAPTER 2 LITERATURE REVIEW

2.1 COMMERCIAL PRODUCTION OF METHANOL

The typical methanol synthesis process consists of three steps: production of syngas,

conversion of syngas to methanol and methanol purification [29].

2.1.1 SYNGAS PRODUCTION

The characteristics of the syngas production depend on the type of feedstock used and

the operation conditions of the process [30], [31].

Natural gas is the primary and preferred feedstock for methanol production in industry

because the corresponding processes for syngas production are low in energy

consumption, capital investment and operating cost [32]. There are three main

technologies applied in industry to produce syngas from natural gas. The most common

approach is an extremely endothermic process called methane steam reforming (MSR),

which is operated at a temperature of 800 -1000 °C and a pressure in the range of 2 – 3

MPa [33]. The reactions are shown in Reaction 2-1 and Reaction 2-2. [9], [34], [35]. The

second common technology for syngas manufacturing is called partial oxidation of

methane (POX) where air is introduced as a source of oxygen. The overall reaction is

exothermic (the sum of Reaction 2-6 to 2-8) [3], [30], [32], [34]. The third well-known

methane to syngas technology is called auto-thermal reforming (ATR) which combines

both steam reforming and partial oxidation into one step. ATR is a thermodynamically

neutral system because heat required for the endothermic steam reforming reaction is

supplied by the exothermic partial oxidation reaction occurring in the same vessel [36].

Reaction 2-1. Methane steam reforming

CH4 + H2O ⇌ CO + 3H2 ∆Ho = 205.43 kJ/mol

Literature review

8

Reaction 2-2. Water gas shift reaction (WGS)

CO + H2O ⇌ CO2 + H2 ∆Ho = −41.00 kJ/mol

Reaction 2-3. Methane partial oxidation

CH4 +1

2O2 ⇌ CO + 2H2 ∆Ho = −35.98 kJ/mol

Reaction 2-4. Oxidation of carbon monoxide

CO +1

2O2 ⇌ CO2 ∆Ho = −282.84 kJ/mol

Reaction 2-5. Oxidation of hydrogen

H2 +1

2O2 ⇌ H2O ∆Ho = −241.42 kJ/mol

Coal is the other main feedstock used for syngas production in countries with low or

costly local availability of natural gas [37]. In this case, syngas is produced via three

processes: gasification, partial oxidation and steam treatment. The reactions are

summarised in Reaction 2-6 to Reaction 2-9 [32], [38].

Reaction 2-6. Oxidation of carbon

C + O2 ⇌ CO2

Reaction 2-7. Reaction with carbon and water

C + H2O ⇌ CO + H2

Reaction 2-8. Water gas shift reaction

CO + H2O ⇌ CO2 + H2

Reaction 2-9. Formation of carbon monoxide via carbon dioxide and carbon

CO2 + C ⇌ 2CO

In recent years, syngas production using shale gas has become popular since large shale

deposits have been discovered in the United States and Canada [39]–[41].

Literature review

9

After possible purification processes, syngas is pressurised with a compressor and added

to recycled syngas and then heated. The produced syngas consisting of H2 and CO with a

ratio of 3 to 5 is introduced to the methanol reactor [42].

2.1.2 METHANOL SYNTHESIS FROM SYNGAS

Methanol was first manufactured by BASF where methanol was synthesised over a

catalyst of ZnO/Cr2O3 at 25 - 35 MPa and 600 - 723 K. The very high pressure required

for the process was undesirable, so that large companies have invested considerable

money and research to improve the operating conditions of the methanol synthesis

process. By 1966, ICI was able to reduce the pressure to 10 MPa by using a new catalyst

based on copper and zinc oxide. Since then this process which is called low-pressure

methanol synthesis and is the only process employed globally for methanol production

[3]. In this process, syngas discharged from the reformer is washed, compressed and

heated before entering the methanol synthesis converter [9]. However, the per-pass

conversion of syngas to methanol is very low (< 20 %), resulting in a large recycle stream

of non-converted syngas. Methanol synthesis is an energy intensive process due to the

high operating pressure and temperature, and a high-volume stream of non-converted

syngas which requires compression to be recycled and mixed with the feed stream

entering the methanol reactor. The reactions that are carried out in the methanol reactor

include [43]:

Reaction 2-10. Methanol synthesis I

CO + 2H2 ⇌ CH3OH ∆Ho = −90.77 kJ/mol

Reaction 2-11. Methanol synthesis II

CO2 + 3H2 ⇌ CH3OH + H2O ∆Ho = −49.58 kJ/mol

Literature review

10

Reaction 2-12. Reverse water gas shift reaction

CO2 + H2 ⇌ CO + H2O ∆Ho = +41.19 kJ/mol

Reaction 2-10 and Reaction 2-11 indicate that the hydrogenation of CO and CO2 are

exothermic and limited by equilibrium reactions. Heat released from the reactions results

in increasing the temperature inside the reactor and hence retards the reaction

conversion. Therefore, in situ heat removal from the methanol synthesis system is crucial

for the reactions to continue and thus for per pass conversion to increase. In methanol

synthesis (Reaction 2-10 and Reaction 2-11), three and four moles are converted into one

and three moles, respectively. Following Le Chatelier principles, in such reactions the

equilibrium will shift the reaction towards the product when high pressure is applied.

Thus, traditionally to achieve a reasonable conversion in the industrial methanol

processes, copper based catalysts and high pressures are mandatory to balance out the

mole decrease during the synthesis [44]. In industry, a variety of technologies have been

engaged to accomplish effective methanol synthesis through gas phase reactions. The

specification and characteristics of each design and technology are categorised based on

flow configuration of feed gas, methods of heat removal and production capacity.

The composition of syngas is characterised by a stoichiometric number S which is shown

in Equation 2-1. Thus, ideally and from a stoichiometric prospective, the best S value for

methanol synthesis is 2. However, a value slightly higher than 2 is preferred for most of

the commercial catalysts [3]. As can be seen from Equation 2-1, a value higher than 2

suggests a high content of H2 while a value less than 2 indicates low concentration of H2.

Equation 2-1. Formula for S value determination

S =moles H2 − moles CO2

moles CO2 + moles CO

Literature review

11

Table 2-1 summarises the current major methanol technology suppliers as well as their

specifically designed reactors and operating conditions [43]. Currently, 61 % of global

methanol production is manufactured via the Johnson Matthey process, 27 % by the Lurgi

process, and 8 % and 3 % by using MGC and Kellog technologies, respectively (Figure 2-1)

[45].

Figure 2-1. Global methanol production technologies distribution

Literature review

12

Table 2-1. The current methanol synthesis suppliers with their methanol convertors/reactors

Reactor systems

Supplier Pressure (MPa) Temperature (°C)

Descriptions Advantages Disadvantages

Single reactors

Johnson Matthey

5 - 10 210-290 A single quench-cooled adiabatic converter with production rate of 3000 tons per day. [46]

Simple and reliable. [47] The heat of reaction is recovered. [3]

The temperature distribution in the reactor is not even. [48]

Haldor-Topsoe 5 - 10 240-270 Isothermal boiling water reactor [49]

Higher reaction conversion and low capital cost. [50] Isothermal boiling water reactor provides a good heat recovery through the production of mid to high pressure steam. [47] Reaction temperature is easy to control. [51]

High operating temperature results in low conversion and hence huge recycle stream[51]. Thermal deactivation of catalysts, thus higher material costs.[47] High pressure drop due to that catalysts packed inside the tube of the reactor. [47]

Mitsubishi 5 - 20 235-270 Double-tubular configuration with methanol catalyst packed in the annular space between the inner and outer tubes

Excellent heat recovery. Steam is produced through the circulation of boiling water. Feed gas is preheated to recover sensible heat. .[52]

Double-tubular configuration lead to higher material costs. Significant pressure drop results in increasing capital costs and imposing difficulties at the reactor operation. [53]

Lurgi 5 -10 230-265 Tube-based reactor with catalysts in fixed tubes. [54]

High conversion and low recycle ratio. [54] High energy efficiency with large amount of steam generation. [3]

High pressure drop in the catalyst packed tubes [55] Large capital expenditure due to the two stage processes. [55]

Linde AG 5 -15 240-270 Isothermal reactor with embedded helically coiled tubes

Tube bundle arrangement inhibits the free fall of the catalysts particles as well as provides higher surface area for heat transfer. [56]

Expensive helical cooling tubes. The energy efficiency is lower than other technologies because it has not implemented sensible heat recovery. [48]

Literature review

13

Reactor systems

Supplier Pressure (MPa) Temperature (°C)

Descriptions Advantages Disadvantages

Multiple reactors

Haldor-Topsoe 4 -12.5 200-310 Collect-Mix-Distribute reactor (CMD reactor). Conventional adiabatic quench reactors. [41]

Extend the lifespan of the catalysts. [47] Higher per pass conversion than conventional adiabatic reactors. [53]

The internal design of the reactor is complex, thus high capital costs of the unit. Large volume of catalysts are required. [48] Difficult to control and optimise the design variables.

Kellogg, Brown and Root (KBR)

(now Halliburton)

N.A. N.A. A series of adiabatic fixed bed reactors with spherical shape. [47]

The thinner wall thickness results in a cheaper manufacturing and installation costs. [50]

Ａ large number of high pressure

reactors, heat exchangers and interconnected piping, also results in high costs. [51]

Toyo Engineering Corporation

N.A. 240-270 The reactor called MRF-Z with production capacity of 5000 – 60000 tons per day Multiple stage radial flow reactors with intermediate cooling using bayonet boiler tubes. [57]

Radial flow of the syngas lead to higher heat transfer coefficient, and less pressure drop.

Literature review

14

2.1.3 SEPARATION AND PURIFICATION

The product stream exiting the methanol reactor is crude methanol since it contains by-

products like water, dissolved gases and traces of H2. The ratio of products is highly

dependent on the type of methanol reactors, feed streams and operating conditions [32].

The purification of methanol and separation of the by-products and impurities are

conducted using one or more distillation columns [32], [58]–[60]. In the current

methanol production plants, two types of distillation trains are applied; cost-saving two-

column distillation and energy-saving three-column distillation [61]. Each type of

methanol distillation column system is suitable for a certain kind of plants. Table 2-2

summarises the advantages and disadvantages of each system as well as the

corresponding schematic diagrams.

Literature review

15

Table 2-2. Selection of methanol distillation trains

Factors Two-column system [59] Three-column system [61]

Selection criterion

The plants with steam reforming When steam import for column reboiler is possible Low capital costs

Insufficient reformed gas waste heat is generated. Large plant capacities Low Energy consumption

Disadvantages High energy consumption High capital costs

Flow diagram [32]

Literature review

16

2.1.4 LIMITATIONS OF COMMERCIAL METHANOL SYNTHESIS PROCESSES

Even though methanol production is well established in industry, the exothermic nature

and thermodynamic characteristics of methanol synthesis reactions present several

difficulties and problems for the optimum process and reactor designs.

HIGH OPERATING TEMPERATURE

The change in free energy (∆G) of methanol synthesis at different temperatures, when P

is set at 3 MPa, is summarised in the Table 2-3 with the values obtained from the

AspenPlus using thermodynamics database in the library. It shows that ∆𝐺 reduces when

temperature decreases, indicating that methanol synthesis reaction is favoured at low

temperature. However, the commercial catalysts (kinetics) shows higher activity at

elevated temperature, therefore the reaction temperature chosen is a compromise of

thermodynamics and kinetics and is within the range of 250 – 300 ℃. Moreover, Figure

2-2 shows that the theoretical CO conversion is limited to approximately 20 % at

commercial operating conditions (300 °C and 50 bar). Higher pressures could improve

the methanol reaction equilibrium and yield higher methanol conversion. However, those

extreme operating conditions result in not only higher investment cost, but also higher

energy demand.

Literature review

17

Figure 2-2. Thermodynamic equilibrium curves for the conversion of syngas to methanol. Syngas with (H2/CO=2) is used

Table 2-3. The change in free energy (∆𝑮) of methanol synthesis at various temperatures

Temperature (°C) The change in free energy, ∆𝐺 (kJ/mol) 25 -167 (l)

100 -149 (l) 200 -83.5 (g) 250 -58.1 (g) 300 -47.0 (g)

DEFICIENT PERFORMANCE OF HEAT TRANSFER

The exothermic nature of the reactions generates a substantial amount of heat, which

requires continuous heat removal to avoid extreme isotherm and maintain isothermality

which help shift the reaction equilibrium to a higher methanol conversion. In addition,

excessive heat and high temperatures influence the catalyst life-span significantly. Heat

distribution within the reactor is not homogeneous due to poor heat transfer. Therefore,

the temperature control within conventional reactors for methanol synthesis is very

difficult.

LARGE RECYCLE RATIOS

Figure 2-3. Equilibrium conversion of CO in methanol synthesis at various temperatures and pressures. Syngas with (H2/CO=2) is used

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18

From Figure 2-3, up to 30 % equilibrium conversion of CO can be achieved at industrial

operating conditions of 250 ℃ and 5 MPa. The equilibrium conversion of CO to methanol

at different operating conditions was calculated by Aspen HYSYS using thermochemical

data within its database. However, the optimum conditions which favour both kinetics

and thermodynamics equilibrium are difficult to balance. As a consequence, CO per-pass

conversion is very low ranging from 6 to 12 %. Thus, a large stream of unreacted high-

pressure gas need to be recycled, i.e., large recycle ratios are required.

Over the past decades, many research groups and institutions around the world have

been working on the development of catalysts, reactor configurations and process

designs that generate methanol at moderate conditions, to increase methanol production

rates thermodynamically. The outcomes can be classified into three approaches: namely

Brookhaven National Laboratory (BNL) homogeneous liquid phase methanol synthesis

at low temperature and pressure (BNL-LTLP) [62], [63], methanol synthesis via methyl

formate as an intermediate (MS via MF), and methanol synthesis using low-grade syngas

at low temperature (MS-LT) [64]–[67]. A summary of these three approaches will be

explained and summarised in the next section. A comparison among these approaches as

well as the commercial methanol synthesis process will be performed, which provides a

potential guideline that helps to determine the research scope and objective.

2.2 METHANOL SYNTHESIS AT MODERATE CONDITIONS

2.2.1 THREE APPROACHES OF METHANOL SYNTHESIS

Approach 1: Brookhaven National Laboratory (BNL) Low Temperature Liquid Phase

methanol synthesis (BNL-LTLP)

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Brookhaven National Laboratory(BNL) proposed a concurrent tandem catalytic system

to produce methanol from syngas with H2/CO ratio of 2 at preferred operating

temperature ranges between 80 - 120 °C and pressure range between 1 - 5 MPa [62], [63],

[68]. The catalytic system consists of NaH, RONa (where R is alkyl group containing 1 - 6

carbon atoms) and M(OAc)2 (where M is Ni, Pd and Co). About 74 % methanol conversion

is obtained at 5.17 MPa and 100 °C. The active catalysts in the system were alkali alkoxide

and Ni(CO)4, which were confirmed by X-ray absorption fine structure (XAFS) analysis

[69], [70]. The reaction rate obtained is five times faster with K alkoxide than Na alkoxide,

and methoxide is a preferred alkoxide ion in the system. Therefore, KOCH3 was selected

as the alkali alkoxide catalyst due to its affordability/availability as well as pronounced

catalysis performance [71]–[73]. The addition of polar solvent, such as

tetrahydrofuran(THF), 1,4-dioxide, triglyme, polyethylene glycol, DMSO, 2-

hydroxybenzothiazole and pyridine, can enhance the methanol conversion and

selectivity [71], [73]–[76]. CO2, dimethyl ether, water and alkali formate were detected

as by-products and the introduction of trace amount of CO2 could deactivate the catalyst

system [71]–[73]. The reaction kinetics using Ni(CO)4/KOCH3 was determined by Wegrzy

et al. and Mahajan et al. over a temperature range between 343.15 K and 393.15 K [73].

By increasing the reaction temperature from 110 to 160 °C, the space time yield (STY)

can also be increased by 5.6 times, resulting in 0.95 kg methanol/L·hr. This STY is higher

than that of ICI commercial methanol production with only 0.5 - 0.77 kg methanol/L·hr

and lower temperature input.

Due to the toxic nature of Ni(CO)4, it has become necessary to find alternatives to avoid

cumbersome handling of the reaction solutions during methanol synthesis. Ni salts,

including NiCl2, NiSO4, Ni(acac)2 and Ni(PPh3)(CO)3 were found to be the good

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20

replacement for Ni(CO)4 for low temperature methanol synthesis, and they were

converted into new active forms that functionalise as catalysts in the process [77], [78].

Mahajan et al. found that 8.9 mol methanol/L catalyst·hr can be achieved when Ni-

salts/base in triglyme/methanol solution was applied at 130 °C and 4 MPa under

continuous 5 % N2 in syngas flow conditions [78]. This result agreed with the commercial

catalysts that yields an equivalent rate of 6 mol methanol/L catalyst·hr at 5 MPa and 250

°C with syngas [79]. In addition to nickel salts, Raney nickel also showed a good catalytic

performance compared with Ni(CO)4. However, the high costs of the catalysts make this

kind of catalysts hard to be applicable in the industry. In addition, since the catalysts are

soluble in the reactants, it is difficult to purify the final products and recycle the catalysts

for further usage.

Approach 2: Low temperature methanol synthesis using low grade syngas (MS-LT)

Low grade syngas is the syngas which contains significant amount of CO2. Tsubaki et al.

postulated a new approach for liquid phase methanol synthesis using the conventional

catalyst of gas phase methanol synthesis - Cu/ZnO/Al2O3 that can tolerate the presence

of traces of CO2 and H2O [66], [67]. The new system effectively showed higher conversion,

and more methanol yield at lower temperature than conventional gas phase methanol

synthesis. The active sites of Cu/ZnO/Al2O3 catalyst for methanol synthesis is not only

metallic Cu but also Cu-Zn site as they both work synergistically to catalyse the methanol

synthesis reactions [80]. Alcohol is required to be present in the low grade syngas system,

and studies showed that the conversion to methanol decreased with increasing carbon

number of alcohol molecules and the secondary alcohols performed the highest activity

with highest conversions [81], [82]. A STY of 0.17 kg methanol/L·hr was obtained when

n-butanol was applied as the promoter at 170 °C and 3 MPa, exhibiting the highest activity.

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A five-step reaction mechanism was proposed based on an in situ diffuse reflection

infrared Fourier transform spectroscopy method, and the pathway is shown in Figure

2-4. [66], [83], [84]. In the mechanism, both CO and CO2 react with adsorbed OH group to

generate acid group, which can further react with alcohol undergoing esterification to

produce the formate group. Finally, the formate group attaches to adsorbed H producing

alcohols as can be seen from Figure 2-4.

Figure 2-4. A proposed mechanism of methanol synthesis using low grade syngas based on the DRIFT study

Approach 3: Methanol synthesis via methyl formate (MS via MF)

Syngas conversion to methanol under mild conditions via methyl formate was first

proposed by Christiansen in 1919 [85]. This route consists of carbonylation of alcohol to

form ester and the hydrogenation reaction of ester to alcohol. The carbonylation reaction

is catalysed by alkali alkoxide and operated at 60 – 120 °C and 1 - 5 MPa (Reaction 2-13).

The hydrogenation of ester to methanol is carried out over copper chromite or copper-

based catalysts at 140 - 180 °C and 1 - 4 MPa (Reaction 2-14).

Reaction 2-13. Carbonylation reaction of alcohol

ROH + CO ⇌ HCOOR ∆H° = −30.97 kJ/mol

Reaction 2-14. Hydrogenation reaction of ester

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HCOOR + H2 ⇌ ROH + CH3OH ∆H° = −52.6 kJ/mol

This approach is characterised by low temperature and liquid phase operating conditions.

The temperature of both reactions is lower than that in traditional gas phase methanol

synthesis. The low operating temperature reduces the thermodynamic limitation and

thus yields higher process conversion. Also, the liquid phase condition allows efficient

heat transfer and hence better temperature control.

Both the two-step process and the concurrent process have been studied in literature

intensively. The reaction kinetics of both reactions and the combined reactions were

proposed using simple power law expressions [65], [72], [86]–[89]. The hydrogenation

reaction has a slower reaction rate than the carbonylation reaction, and it is reported that

the hydrogenation is the rate limiting reaction when two reactions were carried out

simultaneously [72], [86]. The operating temperature ranges are 60 - 100 °C and 120 -

180 °C for the carbonylation reaction and hydrogenation reaction, respectively. Hence a

higher temperature range of 150 to 180 °C was selected for concurrent methanol

synthesis [89]. Alkali formate, CO2 and dimethyl ether were detected in the product of the

carbonylation reaction and CO was identified in the hydrogenation reaction products

[86], [90], [91]. The reaction mechanisms were also proposed for the methanol synthesis

via the methyl formate approach [90].

Comparison of methanol synthesis at moderate operating conditions and the

commercial reactions

Table 2-4 summarises the comparison among the three low-temperature (non-

conventional) methanol synthesis approaches and with the conventional methanol

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synthesis. In summary, all three methods are operated efficiently at moderate conditions

compared to the conventional methanol synthesis and to generate more methanol.

Both the BNL process and MS via MF process require pre-treatment steps to remove

water and carbon dioxide to minimize their negative effects on the catalysts. This

purification step is not practically difficult since advanced adsorption process could

effectively remove CO2 and H2O using zeolites or other adsorption materials. In addition,

regeneration of basic catalysts in the BNL process is compulsory. This can be achieved by

either in situ auto-repair or ex situ regeneration. The toxic nature of nickel catalyst is the

drawback of this process that has inhibited its commercialisation. Therefore,

investigating new catalysts, such as metal-salt catalysts, is essential for the BNL process.

In addition, although the process of low grade methanol synthesis is a good option when

a large amount of CO2 contained syngas is used as feed, additional separation processes

are required for the purification of the final product as large amount of the by-product,

esters, is generated at low partial pressure of hydrogen.

Table 2-4. Comparison of three types of non-conventional methanol synthesis processes with conventional methanol synthesis

Processes Advantages Disadvantages Conventional process

• Mature process with desired economic benefits and industrial outputs.

• At least 50 % of syngas must be recycled.

• High pressure and high temperature process

• High operating costs

BNL process

• Low temperature and pressure, so low operating costs

• High conversion • Low operating costs

• Low tolerance to H2O and CO2, pre-treatment of syngas is required.

• Nickel catalyst is environmentally unfriendly.

• The intermediate, Ni(CO)4 is highly toxic.

• Separation base catalysts with final products.

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Processes Advantages Disadvantages MS via MF concurrent

• It does not require syngas cycle [89].

• Low temperature by applying methanol as the starting material.

• Easier separation of solids catalysts with solvents.

• The reaction rate cannot be increased by using a high space velocity in the combined reactions system [89].

Low grade methanol synthesis

• Low temperature process • No pre-treatment of syngas

is required

• Low conversion compared with the other two processes.

• High pressure is required to ensure relatively high conversions.

• Separation of alcohols in the end is critical if methanol is not used as starting material.

• Higher hydrogen partial pressure is required to consume the intermediate, formate.

• Recycle of the unreacted gas is required.

To apply the methanol synthesis processes at low pressures and low temperatures in

industry, the effectiveness and stability of the catalyst are essential. It is economic

favourable to employ the catalysts that can withstand high pressure and temperature and

can be recycled. Compared with the BNL process and the low-grade methanol synthesis

process, the author has found that the MS via MF is an attractive approach with great

potential. This is especially because the first section of the process, the carbonylation

reaction, is already industrialised and the technology is established. Thus, the focus of

this work is the hydrogenation reaction. Herein, a comprehensive literature review will

be presented in the next section including the reaction rate, roles of catalysts, by-products

formation, and reaction mechanism.

2.2.2 METHANOL SYNTHESIS VIA METHYL FORMATE (MS VIA MF)

CARBONYLATION REACTION

Carbonylation reaction of the MS via MF process for methanol synthesis is defined as the

reaction of carbon monoxide with methanol to yield methyl formate.

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2.2.2.1.1 Reaction mechanism

Christiansen postulated the mechanism of carbonylation reaction for a system including

alcohol and alkali alkoxide. The mechanism suggests that the alcoholate ion (RO-) donates

electrons to the unused 2p orbitals of carbon atoms forming a complex (Reaction 2-15),

which subsequently reacts with alcohol to obtain the active catalyst (Reaction 2-16).

Tonner et al. determined that Reaction 2-15 is the rate-limiting step [90].

Reaction 2-15. Mechanism step 1

RO:−+ C = O ⇌ (RO − C = O)−

Reaction 2-16. Mechanism step 2

(RO − C = O)− + ROH ⇌ ROOCH + RO:−

2.2.2.1.2 Reaction kinetics

The carbonylation reaction is a first order reaction with respect to carbon monoxide, as

found in all reaction kinetics studies. Since the alcohol feed is always in large excess, it is

not possible to determine the reaction order with respect to alcohol. Tonner et al. has

performed a kinetics study of the carbonylation reaction using various types of alcohols

ranging from C1-C4 (R-OH). The system was catalysed by sodium alkoxide (Na-OR),

which was synthesised in situ from sodium metal dissolving into alcohol [72]. The results

of the study indicated that secondary butanol showed the highest reaction rate, while

methanol corresponded to the slowest. This suggests that the electron-directing effect of

alcohol might be an effective factor that controls the reaction rates. Also, the increase in

the length of the alcohol chain and/or the degree of substitution next to the hydroxyl

group result in enhancing the rate constant. Further, by looking at the reaction

mechanism, any substituent group that enhances the electron density on an O atom will

accelerate the reaction if Reaction 2-15 affects the overall reaction.

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Liu et al. studied and developed a kinetics expression (Equation 2-2) for the

carbonylation reaction over KOCH3 catalyst over a temperature range of 60 – 110 °C and

2.5 - 6.5 MPa [86]. A power-law expression was used to determine forward and reverse

reactions in the carbonylation reaction. Two methods are commonly used to determine

reverse reaction; the first one is by fitting a pressure-time plot and the second one is by

using equilibrium results at the end of experiment. However, a large discrepancy was

observed between the two methods when the concentration of methyl formate was high.

The reason is that, in the first method the rate of reverse reaction was obtained at low

concentration of methyl formate, while in the second method it was obtained when the

system was at equilibrium condition and methyl formate concentration was much higher

than in the initial stage. In addition, the reaction rate differences were due to changes in

the activity coefficient and CO solubility between high and low methyl formate

concentration.

Equation 2-2. The overall reaction rate expression [86]

r1 = 2.88 × 109e−10126

T CCAT1CMEOHPCO − 1.19 × 1019e−16788

T CCAT1Cmef

Another study of carbonylation reaction kinetics was performed by Liang et al. for the

NaOCH3/CO system at a temperature range of 60 – 110 °C and pressure ranges between

2 -4 MPa. The kinetics of both forward and reverse reactions were determined [87]. By

assuming pressure drop was only due to the disappearance of CO to form the desired

product, methyl formate, the forward reaction rate was derived. In addition, the

equilibrium constant, Ke, was used to generate the kinetics expression of the reverse

reaction. The overall reaction rate for the carbonylation reaction using NaOCH3 can be

found in Equation 2-3 which can be applied in the range of 60 – 110 °C. The activation

energy of the forward reaction is 70.7 kJ/mol. However, using the equilibrium constant

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to derive the reverse reaction rate is questionable as the carbonylation reaction is not an

elementary reaction.

Equation 2-3. The overall reaction rate expression [87]

𝑟 = k−1CcatCmCCO − k−1CcatCmf

= 1.414 × 109 exp (−70748

RT) CcatCmCCO

− 2.507 × 1012 exp (−92059

RT) CcatCmf

In another study conducted by Chen et al., the kinetics expressions of the forward

reaction was determined with and without a promoter as shown in Reaction 2-4 and

Reaction 2-5 [88].

Equation 2-4. Carbonylation reaction forward reaction rate [88]

−r = k1p(CO) = 9.96 × 106 exp (−67630

𝑅𝑇) p(CO)

Equation 2-5. Carbonylation reaction forward reaction rate with pyridine as a promoter [88]

−r = k1p(CO) = 8.82 × 106 exp (−61190

𝑅𝑇) p(CO)

Tonner et al., on the other hand, believed that limited solubility of CO in alcohol solvents

was not responsible for the reaction rates since the highest solubility of CO in methanol

showed the slowest reaction rate. Moreover, the reaction kinetics ware independent of

catalyst concentration if adequate catalyst was present, this can be explained by a high

amount of alcoholate ions (RO-) in the solution increasing the rate-limiting reaction

(Reaction 2-15).

2.2.2.1.3 The role of catalysts

The effects of using ethanol with different alkali ethoxide catalysts on the reaction rate

was studied [90]. It was found that the rate of the carbonylation reaction increased with

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the order of K>Na>Li. As Reaction 2-15 determines the overall carbonylation reaction,

the formation of alcoholate ion (RO-) is dependent on the ionization potential of the alkali

metal. The ionization potential of metal alkali K, Na and Li, are 4.31 V, 5.12 V and 5.36 V

respectively. Hence, a fast carbonylation reaction rate was observed when potassium

ethoxide was used. In addition, higher reaction rates were obtained using

KOCH3/methanol compared to NaOCH3/methanol system, which indirectly indicates the

higher catalytic activity of KOCH3 and the important role of alcoholate ions in the reaction

as discussed earlier [86].

2.2.2.1.4 Role of promoters

In order to improve the carbonylation reaction rate, several solvents were added to the

reaction system and their effects were investigated. It was found that in the presence of

a solvent, or as named a promoter, like pyridine, the carbonylation reaction rate was

enhanced more than 1.5 times that in the absence of the promoter [88], [92]. No

explanation was given, and no fundamental studies have been conducted to identify the

reasons behind the positive effects of the promotors. However, adding promoters could

increase the subsequent difficulty in product separation.

2.2.2.1.5 By-products formation

Traces of alkali formate was detected in the product of the carbonylation reaction as

proposed by Tonner et al. and Liu et al. [72], [86]. Possible reactions, such as Reaction

2-17 or Reaction 2-18 might result in the formation of alkali formate [90]. Also, traces

amount of CO2, and dimethyl ether were detected as by-products. The by-product

formation increased with reaction temperature and reaction time [86].

Reaction 2-17. Formation of sodium formate [90]

HCOOR + RONa ⇌ R2O + HCOONa

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Reaction 2-18. Formation of sodium formate [90]

NaOH + HCOOR ⇌ HCOONa + ROH

The effect of CO2 on the carbonylation reaction was investigated at 70 °C. No reaction was

observed after purging CO2 into the system. It is evident that CO2 has a significant

consequence on the carbonylation reaction, and the probable reason could be

deactivation of the catalysts by CO2 to generate CH3OCOOK as shown in Reaction 2-19.

Reaction 2-19. Catalyst deactivation by CO2

CO2 + CH3OK ⇌ CH3OCOOK

HYDROGENATION REACTION

Hydrogenation or sometimes called the hydrogenolysis reaction, is the second stage of

methanol synthesis via the methyl formate approach. In this reaction methyl formate

reacts with H2 to form methanol. Various heterogeneous copper-based catalysts have

been tested in both gas phase and liquid phase systems in a slurry reactor. The following

discussion covers this reaction.

2.2.2.2.1 Reaction kinetics

The kinetics expression of the hydrogenation reaction was performed by Monti et al.

(Equation 2-6) with regards to the concentration of methyl formate, hydrogen pressure

and catalysts concentration in the temperature range of 408 and 473 K. The activation

energy was determined to be 62.4 ± 0.2 kJ/mol at a hydrogen pressure of 4.5 MPa [65]. A

modified kinetic expression for the hydrogenation reaction was determined in the

presence of CO (P > 80 kPa) and at a temperature of 446 K (Equation 2-7).

In addition, the effects of mass transfer (controlled by agitation speed) on reaction rate

was determined. It was found that the reaction rate was not subject to mass transfer

limitations when the catalyst concentration was above 20 g/L [65].

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Liu et al. found that the hydrogenation reaction was a highly selective reaction with no

by-products [86]. A reaction rate expression was derived in the range of 120 to 200 °C by

fitting experimental data to kinetic models using a non-linear regression method

(Equation 2-8). The reverse reaction was not included in the expression because the

equilibrium conversion was very high. As seen from the denominator of the equation, the

square root term of methyl formate indicates that methyl formate dissociated after being

adsorbed on the copper-chromite catalyst has effects on the reaction rates, while the

absence of a dependent on H2 suggests that adsorption of H2 on the catalyst surface is

insignificant. The presence of the CO term suggests that the inhibitory effect of CO on

reaction rate is because of competitive adsorption of CO and methyl formate on the same

active sites.

Equation 2-6. Reaction kinetics expression for the hydrogenation reaction [65]

−rA = 8.35 × 103e−7510

T CmfPH2Ccat

Equation 2-7. Reaction kinetics expression for the hydrogenation reaction in the presence of CO [65]

−rA = 7400 e−7510

T CmfPH2CcatPCO

−0.32

Equation 2-8. Reaction kinetics expression for the hydrogenation reaction [86]

𝑟𝐴 =1871.5e−

69400RT CmefPH2

Ccat

1 + (0.039Cmef)12 + 0.096PCO

2.2.2.2.2 The role of catalysts

Copper-chromite is used to catalyse the hydrogenation reaction after complete reduction

under H2 or N2/H2 atmosphere. In the reduction process, copper chromite (CuO, CuCr2O4)

is reduced to Cu0 and CuCrO2 [93]. Different reduction conditions, such as heating time

and heating temperature, were used by various authors. The BET surface area and copper

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surface area of G-89 (Nissan Girdler Catalysts Company, Ltd) copper chromite after

reduction was reported to be 24.8 m2/g and 4.9 m2/g, respectively [65].

2.2.2.2.3 By-products formation

It has been reported that in gas phase hydrogenation reactions, the undesired

decarbonylation reaction can take place, producing CO as a by-product (Reaction 2-20).

The production of CO exhibits a negative effect on the system as it inhibits the

hydrogenation reaction by replacing the adsorbed hydrogen on the catalyst surface and

causes catalytic deactivation especially when the hydrogenation reaction is carried out in

the gas phase at atmospheric pressure [91].

Reaction 2-20. De-carbonylation reaction

HCOOCH3 ⇌ CH3OH + CO

In order to investigate the influence of CO on the liquid-phase hydrogenation reaction, a

number of experiments were carried out with different concentrations of CO in the

system [65]. It was found that the effects of the added CO is inhibition of kinetics rather

than progressive catalyst poisoning or thermodynamic equilibrium restriction.

Therefore, poisoning of copper chromite catalysts was not the case in liquid phase

hydrogenation reaction. However, Liu et al. observed that the existence of CO reduced the

hydrogenation reaction rate, while such deleterious effects can be partially restored

when CO was removed [86]. In addition to CO, CO2 was also found to be a factor that

caused catalyst deactivation and such an effect cannot be reversed [86] and CO2 is

thought to poison the catalyst.

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CONCURRENT LOW TEMPERATURE METHANOL SYNTHESIS (TWO-REACTIONS

IN ONE REACTOR)

Combining the carbonylation reaction with the hydrogenation reaction into one reactor

has been studied in the literature [91], [93]–[95]. Since the hydrogenation reaction has a

much slower reaction rate than the carbonylation reaction, Liu et al. reported that the

hydrogenation reaction was the rate limiting reaction when two reactions were carried

out simultaneously. In addition, it was found that methanol production in the concurrent

system was significantly higher than in the individual reaction system [89]. Liu et al.

believed that certain synergetic effects might contribute to this phenomenon, but no

further discussion was conducted.

2.2.2.3.1 Operating conditions

In the concurrent low temperature methanol system, H2/CO ratio, temperature, and

pressure are the three major dominating conditions for liquid phase methanol synthesis.

Based on the thermodynamic study conducted by Liu et al., when the feed stream H2/CO

ratio was maintained at 2, the possible concurrent methanol synthesis reactions can only

proceed in the operating ranges within the shaded area shown in Figure 2-5 [89].

Liu et al. has combined the kinetics expressions of the individual reactions to predict the

rate for the concurrent system of methanol synthesis, as shown in Equation 2-9. It was

assumed that the hydrogenation was a rate-limiting reaction and the carbonylation

reaction of methanol was in equilibrium. The optimum operating conditions were

determined by taking the derivative of Equation 2-9 with respect to PH2/PCO. It was found

that the optimum operating temperature and pressure was between 140 to 180 °C and 3

– 6 MPa, respectively.

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Figure 2-5. Thermodynamic analysis of methanol reaction condition when H2/CO = 2 [89]

Equation 2-9. Predicted concurrent methanol production [89]

r =6.31 × 10−4e−

4568T Ccat2CmeOHPCOPH2

1 + 0.096PCO

A higher temperature condition was preferred in the concurrent methanol synthesis

process as indicated by Liu et al., because higher reaction rates can be obtained when

higher temperature is applied. However, in practice, at 180 °C, the methanol production

rate was lower than that at 140 and 160 °C. This is probably due to the catalyst

deactivation by CH4, which was the only by-product observed at 180 °C [89]. The reaction

rate increases at higher pressure, whereas methanol production decreases due to catalyst

deactivation.

As CO is a major reactant in the carbonylation reaction and similarly H2 in the

hydrogenation reaction, it is important to control the H2/CO ratio. It was found that at

higher CO concentration (low H2/CO ratio) more methyl formate was formed from

methanol carbonylation. The produced methyl formate reacts with H2 producing

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methanol, therefore, the production rate increases. However, with excess CO, catalyst

deactivation is inevitable. The optimal ratio, on the other hand, cannot be determined

because of the effect of the other reaction factors including temperature, pressure and

catalysts loading.

2.2.2.3.2 Combination of catalysts

Palekar et al. investigated the concurrent methanol system experimentally at 150 °C and

5 MPa [93]. The alkali methoxide and copper chromite, which were used to catalyze the

carbonylation reaction and hydrogenation reaction were combined and fed into the

reactor [93]. Between the two reactions, the hydrogenation reaction is the rate-limiting

reaction because its rate is two orders of magnitude lower than the carbonylation

reaction [86]. The results showed that methanol production rate was higher than the

hydrogenation reaction alone, which suggested that a synergistic behaviour of the two

reactions was taking place [93]. In addition, better tolerance to the presence of CO2 and

H2O was observed in the concurrent systems when the operating temperature was above

100 °C [89], [93].

Since the catalyst, KOCH3, used solely for the carbonylation reaction deactivates

gradually, it is suggested that the addition of the copper chromite contributes to

regeneration of the deactivated catalyst. A proposed mechanism of the catalyst

regeneration was put forward and further proved by several experiments using alkali

carbonate, formate, hydroxide and bicarbonates with copper chromite [88]. Therefore, a

synergistic effect between KOCH3 and copper-chromite leads to a higher rate of methanol

production in the concurrent synthesis [93], [95].

In a study conducted by Liu et al., it was found that methanol production decreased with

increasing the concentration of potassium methoxide. This is probably due to the

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blockage of the active sites on copper chromite surface by the alkali ions and methoxide

ions block, and this phenomenon is so called alkali site blocking [95]. Thus, different alkali

methoxides were used with copper chromite to determine the effects of alkali site

blocking. It was found that potassium methoxide provided the highest methanol

synthesis rates, however, the reasons were not clearly identified. The conversion rate of

methyl formate increases from sodium to cesium, or in another words, with decreasing

the ionization potential of alkali metals [72]. With regard to rubidium methoxides and

cesium methoxides, the methanol production rate is very small. This is probably due to

the large size of rubidium and cesium ions that occupy more hydrogenation sites on the

catalyst surface decreasing the methanol production rate. Alkaline earth methoxides can

also be used as catalysts for methanol synthesis, however they showed lower methanol

production rate than that of alkali methoxides [95].

There is a conflict in regard to identification of copper chromite active sites. Palekar et al.

believed that it was likely that Cu+ is the active species in the concurrent synthesis [93].

However, there are other research groups who propose that Cu2+, Cu0 on CuCr2O4, Cu0 on

or mixed with Cr2O3 , Cu2O or CuO on CuCr2O4 or Cu0 on Cu2Cr2O4 can be the active species

[93].

Ohyama compared the performance of several copper-based catalysts along with KOCH3.

His results indicated that the commercial catalysts N203SD and G-89, and copper

chromite showed activity for methanol synthesis, particularly, N203SD which exhibited

the highest space time yield (STY) [96]. CuO/ZnO based catalysts including KMB and

G668 showed no catalytic activity for the hydrogenation reaction, which agrees with the

results shown by Palekar et al. [95]. In addition, methanol is only produced when both

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CuO and Cr2O3 are employed, suggesting that the active site for hydrogenation reaction

might be obtained by mixing CuO and Cr2O3.

2.2.2.3.3 High selectivity for methanol over DME, water and CO2

Small numbers of by-products were observed and identified such as CO2, DME and H2O,

which were the same by-products observed with the carbonylation reaction [93]–[95],

[97]. The concentration of by-products increased at higher temperature. Methane, for

example, was detected as by-product when temperature was increased to 180 °C. In the

concurrent reaction system, the overall reaction rate was observed to be considerably

lower than predicted. This was attributed to the presence of methane which might

deactivate the catalyst [89]. Although the negative effect of CO2 on concurrent methanol

production has been observed, the addition of CO2 can significantly reduce the formation

of by-products, i.e. DME formation [98].

ADVANTAGES OF THE CON-CURRENT METHANOL SYNTHESIS COMPARED TO

THE INDIVIDUAL REACTIONS

Liu et al. summarised two major advantages with bringing the two reactions together in

one reactor [98]. These advantages include higher methanol production and enhanced

tolerance to CO2. At 140 and 160 °C, the concurrent synthesis (two reactions at once) gave

80 and 50 % higher yield respectively than the separate carbonylation and hydrogenation

[98]. It was found that the non-reversible poisoning effect of CO2 on CH3OK had

significantly reduced and become reversible in the con-current methanol synthesis. Liu

et al. attributed this synergistic effect to the interaction between the two catalysts in the

system. The homogeneous catalysts (CH3OK) and its ionic form (CH3O-) are adsorbed on

the copper chromite surface and compete for the same active sites as CO, CO2 and H2O,

which results in a higher methanol reaction rate.

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2.3 SOLUBILITY IN GAS-LIQUID PHASE

The reactants involved in the current research are gas and liquid phases. Since the

reaction takes place in the liquid phase (where the catalyst is), the gas reactant must first

be contacted with the liquid and dissolve the liquid reactant. During the solid catalytic

process, gas and liquid reactants must diffuse or move to the catalyst surface to trigger

the reactions.

Therefore, the solubility of the reactants will affect the movement from phase to phase,

which means the reaction rate can be influenced by mass transfer. To understand gas-

liquid solubility at the typical reaction operating conditions, a short review will be

provided based on the fundamental experimental and modelling works for the proposed

fluid systems.

2.3.1 DEFINITION OF SOLUBILITY

According to the official IUPAC nomenclature, the definition of solubility is: the analytical

composition of a saturated solution expressed as a proportion of a designated solute in a

designated solvent.

The term ‘saturated’ refers to equilibrium of vaporization and dissolution of solute in the

solvent like CO and/or H2 gas in methanol and/or methyl formate liquid. In general, we

are concerned with the liquid mixture which at temperature and pressure is in

equilibrium with a vapour mixture at the same conditions. Thus, if a gaseous phase and a

liquid phase are in equilibrium, then for component i in the mixture, the condition of

thermodynamic equilibrium is given by Equation 2-10.

Equation 2-10.Thermodynamic equilibrium of the fugacity in gas phase and liquid phase

fiV = fi

L

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In the following discussion, we will only focus on the pure gas in pure liquids and neglect

effects due to the surface tension, gravitation, electric or magnetic fields, semipermeable

membranes, or any other special conditions.

2.3.2 FUGACITY OF THE SYSTEM

The fugacity of a component in a mixture depends on the temperature, pressure and

composition of the mixture.

With respect to the vapour phase, the composition is expressed by the mole fraction y.

The fugacity of i in the vapour 𝑓𝑖𝑉 can be related to temperature, pressure and mole

fraction by the fugacity coefficient, ϕi is expressed by Equation 2-11.

Equation 2-11. The equation of the fugacity coefficient

ϕi =fi

V

yiP

The fugacity coefficient, ϕi, is normalised by the partial pressure of component i, which

in turn, as P approaches to 0, ϕi tends to 1. Therefore, at low pressure, it is usually a good

assumption to set ϕi = 1. The term ‘low’ depends on the mixture system. For typical

mixtures of nonpolar (or slightly polar) fluids at a temperature near or above the normal

boiling temperature of the least volatile component, ‘low’ pressure means a pressure less

than a few bars. For mixtures containing one component of very low volatility and

another of high volatility, the fugacity coefficient of the light component may be close to

unity for pressures up to 1 – 2 MPa.

Regarding the liquid phase, the fugacity of component i in liquid phase is related to the

composition of that phase through the activity coefficient, γi . Like the vapour phase

situation, the mole fraction of component i in the liquid phase, x, is usually used to

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describe the fugacity of components in that phase. The relationship between fugacity in

the liquid phase and mole fraction can be determined by Equation 2-12.

Equation 2-12. The equation of activity coefficient

γi =ai

xi=

fiL

xifi0

Where 𝑓𝑖0 is the standard-state fugacity of component i (sometimes, it refers to the

reference state) and ai is the activity of component i. For most cases, activity coefficients

for solutions contained nonelectrolytes are based on a standard state, for every

component i, 𝑓𝑖0 is the fugacity of pure liquid i at system temperature and pressure, that

is to say, the pressure is the total pressure and the composition is 1.

2.3.3 SOLUBILITY LITERATURE DATA REVIEW

The system of liquid phase methanol synthesis involves two types of reactions;

carbonylation of methanol to methyl formate and hydrogenation of methyl formate to

methanol (Reaction 2-21 and Reaction 2-22).

Reaction 2-21. Carbonylation reaction

CH3OH + CO ⇌ CH3OCOH

Reaction 2-22. Hydrogenation reaction

CH3OCOH + 2H2 ⇌ 2CH3OH

Initially CO gas is fed to a reactor containing liquid methanol at a desired temperature

and pressure. Every one mole of CO reacts with one mole of methanol generating one

mole of methyl formate. Thereafter H2 is fed to the system where one mole of methyl

formate reacts with two moles of H2 producing two moles of methanol. Therefore, to

investigate those reactions and comprehend the role of feed gas solubility in liquid phase

components in the system, four major liquid-gas systems need to be studied including

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CO-CH3OH, H2-CH3OH, CO-CH3OCOH, H2-CH3OCOH. Since substantial work has been

carried out based on systems where hydrogen gas was involved, we introduced this first

HYDROGEN SOLUBILITY

2.3.3.1.1 Solubility of Hydrogen in methanol

The published vapour-liquid solubility data for the hydrogen-methanol binary system is

mostly over a temperature range from 298 to 343 K and a pressure range between 0.5 to

4 MPa [99]–[102]. Liu et al. extended the temperature range from 343 to 413 K in order

to satisfy their study requirements [103]. In addition to that, Bezanehtak et al. and

Francesconi et al. investigated the solubility of H2 in methanol at moderate pressures up

to 11 MPa [104], [105]. Higher pressure ranges from 5 to 100 MPa have been intensively

studied by Brunner et al. [106]. Furthermore, a few experiments were performed at

temperatures lower than 298 K for various pressure ranges [107]–[109]. Table 2-5

summarises the temperatures and pressure ranges for which H2-methanol solubility data

is available.

Table 2-5. Published literature for H2-CH3OH binary system

Source Temperature (K) Pressure (MPa)(1)

Radhakrishnan et al. [99] 298 – 343 0.1 Choudary et al. [100] 293, 308, 318, 328 0.463 – 2.1 Wainwright et al. [101] 291 1.1 – 3.62 Liu et al. [103] 293 – 413 0.5-1.6 Bezanehtak et al. [104] 278.15, 288.15, 298.15, 308.15 2 – 11 Brunner et al. [106] 298.15, 323.15, 373.15 5.08 – 110 Francesconi et al. [105] 323.8 - 476.6 4.78 – 10.98 Katayama et al. [107] 213 – 298 0.1 Luhring et al. [102] 293.2 0.1 Descamps et al. [108] 248.41 – 308.20 0 – 3 Gemo et al. [109] 278 0.37 – 1.56

(1) Pressure is either total pressure or partial pressure of the hydrogen since they are not clearly indicated or clarified in most publications

There are several methods applied in the literature to determine the solubility of H2 in

methanol. One of the common methods is the static-analytic technique with liquid phase

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recirculation and online gas chromatography where the sampling was achieved through

rotating sampling valves [101], [104], [108]. However, the sampling valve for liquid

recirculation line poses a problem of introducing some of the carrier gas when switching

the valve between vapour and liquid samples. Also, the sampling valve for the vapour

products requires a great deal of attention to eliminate the pressure drop and obtain

reliable samples.

Liu et al. and Gemo et al. utilized a synthetic method with total pressure measurement at

the desired temperature (figures can be found in Table 2-5). In this method, H2 is injected

into an autoclave loaded with methanol until the pressure inside the autoclave reaches

the desired level at which point stirring starts. The amount of absorbed H2 was

determined by the pressure change between pressure before and after stirring [103],

[109]. The ideal gas law was applied to evaluate the number of moles of both solutes and

solvents. The drawback of this method is the uncertainty of the amount of H2 absorbed in

the methanol before stirring. The quantity determined in such a methodology would be

usually underestimated.

Francesconi et al. applied a similar method to measure hydrogen solubility in alcohols

[105]. However, in this work a more sophisticated procedure is used to calculate the

number of H2 moles in the liquid phase. More details are provided in the modelling

section (section 2.3.4). Descamps et al. applied both synthetic methods and static-analytic

method to the H2-methanol binary system at 298 K to compare the outcomes of the two

methods and to identify the consistency of those two techniques. They have found that

the results from both methods are similar, which means both methods can be used

reliably for in the solubility measurements [108].

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The method employed by Radhakrishnan et al. and Choudary et al. is to saturate methanol

with hydrogen at a desired temperature and pressure in an autoclave [99], [100]. After

saturation, a known amount of saturated liquid sample was transferred to a burette,

where the dissolved hydrogen gas was desorbed/released to the atmosphere. The

volume of desorbed hydrogen can be determined at atmospheric pressure. The drawback

of this method is only single point at specific temperature and pressure can be

determined in every experiment to ensure the accuracy of the results. Katayama et al. also

determined the volume of solvent and vapour phase to measure the solubility of H2 in

methanol by using a static method with mercury displacement [107]. Nevertheless, such

a method can be used for the determination of gas solubility at low temperature due to

the limitations of the equipment.

Figure 2-6 and Figure 2-7 display P-x diagrams of available literature solubility data at

approximately 295 K and a pressure range from 0.5 to 100 MPa. To ensure the unit

consistency, some data points in Figure 2-6 have been converted from the raw literature

data. As can be seen in Figure 2-6, the solubility data of H2 in methanol is in a good

agreement for many researchers except Choudhary’s work, even though different

experimental procedures and methodology were applied. Since the method applied by

Choudhary et al. suffers of limitations in determining the solubility, no other solubility

work performed after 1986 used their procedure. At high pressure, the results from both

Bezanehtak et al. and Brunner et al. are almost overlapping, which provides a good

guideline for the current research. The results indicate that with pressure increase, the

solubility of hydrogen in methanol increases linearly as expected, in accordance with

Henry’s law.

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At high pressures, the non-ideality of gas is profound due to the compressibility of gas at

high pressures. Therefore, the fugacity is required to use correct the deviation at high

pressures. Based on the Krichevsky-Kasarnovsky equation (Equation 2-13), the deviation

of Henry’s constant can be described at high pressures, and the solubility of gas becomes

smaller than that if only Henry’s law is applied. Hence, as shown in Figure 2-7, the linear

correlation of pressure and mole fraction is no longer showed at high temperature ranges.

Equation 2-13. Krichevsky-Kasarnovsky equation

ln (f2

x2) = ln(H2,1) +

v2∞(P − P1

s)

RT

0.000 0.002 0.004 0.006 0.0080.0

0.5

1.0

1.5

2.0

2.5

3.0

3.5

4.0

Pre

ssure

(M

Pa

)

Mole fraction

Figure 2-6. Literature data of H2-methanol binary system at room temperature at low pressures. The x-axis is mole fraction of H2. The y-axis is total pressure. (+) Choudhary et al. at 293 K [100]; (*) Wainwright et al. at 291 K [101]; (∆) Descamps et al. at 291.2 K [108]; (○) Liu et al. at 296.25 K [103]

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0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.140

20

40

60

80

100

Pre

ssure

(M

Pa

)

Mole fraction

Figure 2-7. Literature data of H2-methanol binary system at room temperature at high pressures. The x-axis is mole fraction of H2. The y-axis is total pressure. (squares) Brunner et al. at 298.15 K [106]; (solid circles) Bezanehtak et al. at 298.15 K [104]

In addition, the published data of hydrogen solubility in methanol at various

temperatures are displayed in Figure 2-8 and Figure 2-9. (Due to different axis-scale,

separate graphs were shown). The results clearly show that the solubility of hydrogen in

methanol increases with temperature increase which is different to other gases, in which

solubility usually decreases with temperature increase.

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0.000 0.001 0.002 0.003 0.004 0.005 0.006 0.0070.0

0.5

1.0

1.5

2.0

2.5

3.0

3.5

4.0

373.95 K

363.55 K

308.2 K

Pre

ssure

(M

pa)

Mole fraction

Figure 2-8. Literature data of H2-methanol binary system at various temperature at low pressure. (X) Liu et al. at 373.95 K [103]; (solid square) Liu et al. at 363.55 K [103];(Φ) Descamps et al. at 308.2 K [108]

0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.14 0.16 0.180

20

40

60

80

100

120

298.15 K

323.15 K

373.15 K

Pre

ssure

(M

Pa

)

Mole fraction

Figure 2-9. Literature data of H2-methanol binary system at various temperature at high pressure [106]

In this work of low pressure and low temperature methanol synthesis, it is crucial to

understand hydrogen solubility in methanol at temperatures in the range of 50 – 100 °C

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and for pressures in the range of 1 to 3 MPa which are the experimental conditions for

the hydrogenation reaction. However, the available literature data are not sufficient to

cover this range. Therefore, we have performed the necessary solubility experiments of

H2 in methanol at our desired operating conditions.

2.3.3.1.2 Solubility of Hydrogen in methyl formate

To date, only Liu et al. and Wainwright et al. have performed experiments on H2 solubility

in methyl formate, and both of them aimed to study the reaction kinetics on

hydrogenation of methyl formate [101], [103]. As discussed in section 2.3.3.1.1, Liu et al.

applied a synthetic method with total pressure measurement and Wainwright used a

static-analytic technique by measuring the liquid and gas sample composition via a GC.

Wainwright’s work only focused on one temperature which is 291 K, whereas Liu studied

a wider range of pressures and temperatures.

As can be seen from Figure 2-10 (a), the data of Liu et al. are plotted as six isotherms with

three data points have been measured. However, there are several issues about this

system. Firstly, the isotherms do not pass through the origin except the isotherm at

313.85 K. Secondly, the linearity of the isotherm was not confirmed which contradicts

Henry’s law where partial pressure is proportional to the mole fraction. Furthermore,

there was no discussion in the paper, to indicate what if any uncertainties have occurred

in the measurements. The authors applied the ideal gas law to calculate the mole

fractions, which might be not applicable at high pressures and temperatures. Besides,

methyl formate is highly volatile, it is essential to exclude the partial pressure of methyl

formate from the total pressure when determining the mole fraction at high temperature.

Because no other experiments have been conducted on the H2-methyl formate system at

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similar operating conditions as Liu et al., it is difficult to comment on the accuracy of their

results.

Figure 2-10 (b), on the other hand, which presents Wainwright et. al.’s results shows that

they have employed two different methods in determining the solubility of hydrogen in

methyl formate. In contrast to the H2-methanol system, the results of the two methods

were inconsistent due to the experimental apparatus and procedures being sensitive to

the presence of a highly volatile compound, in this case methyl formate. No further

comparison or explanation was proposed by the author regarding the discrepancy

between the two methods. However, this author believes that their apparatus is

sufficiently accurate to measure the solubility of hydrogen in methyl formate since the

other results obtained from their instrument is in a good agreement with other work. Like

Liu’s work, the lack of solubility data in literature at conditions similar to theirs resulted

in an absence of a comparison factor for accuracy confirmation.

0.000 0.001 0.002 0.003 0.004 0.005 0.006 0.0070.0

0.2

0.4

0.6

0.8

1.0

1.2

1.4

1.6

1.8

(a)

313.85 K

323.05 K

345.35 K

353.95 K

363.85 K

373.45 K

Pre

ssure

(M

Pa)

mole fraction

0.000 0.001 0.002 0.003 0.004 0.005 0.006 0.0070.0

0.5

1.0

1.5

2.0

2.5

3.0

3.5

4.0

4.5

5.0

(b)

method by Wainwright et al.

method by Albal et al.

Pre

ssure

(M

Pa)

Mole fraction

Figure 2-10. Solubility data of H2-methyl formate binary system. The x-axis is the mole fraction of H2. The y-axis is the total pressure. (a) by Liu et al. [103]; (b) by Wainwright et al. [101]

This limitation of literature data for H2-methyl formate binary system motivated the

author of this work to run further experiments to determine the solubility of hydrogen in

pure methyl formate.

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CARBON MONOXIDE SOLUBILITY

Conventionally, the carbonylation reaction is operated over a temperature range 50 – 100

°C and pressures above 1 MPa. The carbonylation reaction system consists of methanol

and CO, and after the reaction commences, methyl formate is formed. Thus, a literature

review on the solubility of CO-methanol and CO-methyl formate systems is necessary. To

the best of the author’s knowledge, no work has been published on the CO-methyl

formate system. Hence, the following review covers only the solubility of CO in methanol.

2.3.3.2.1 Solubility of Carbon monoxide in methanol

Due to the high toxicity of carbon monoxide (CO), only a few studies have been dedicated

to investigating CO solubility in methanol or any other solvents. Tonner et al. and Liu et

al. examined the solubility of carbon monoxide as a requirement to study the kinetics of

the carbonylation reaction [110], [111]. Tonner studied CO solubility at temperatures of

298 K and 323 K and pressures up to 4 MPa using a static-analytic technique through gas

and liquid sampling using a GC, while Liu performed experiments over a wider

temperature range of 293 K to 413 K and pressures between 0.5 to 1.7 MPa using the

synthetic method with total pressure measurement. A series of solubility data at a higher

pressure range were measured by Brunner et al. [106]. In another study conducted by

Luring et al., CO-methanol solubility results were obtained in terms of Henry’s constant

at a temperature of 293.2 K [102].

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0.000 0.005 0.010 0.015 0.020 0.025 0.0300

1

2

3

4

5

6

7

8

9

Pre

ssure

(M

Pa

)

Mole fraction

Figure 2-11. Solubility data of CO-methanol binary system at 323 K. The x-axis is the mole fraction of H2. The y-axis is the total pressure. (x) Liu et al. [111]; (○)Tonner et al.[110]; (∆) Brunner et al. [106]

Figure 2-11 shows that the isotherms from Liu et al. and Brunner et al. studies are in a

reasonably good agreement even though different methodologies were applied. Tonner’s

work shows a discrepancy to Brunner’s results. However, both work show comparable

results of H2-methanol system which indicate that the different results are not strongly

affected strongly by the measurement apparatus. This is possibly because methanol is not

volatile compare to methyl formate, and it does not require a highly sophisticated

instrument to compute the moles of carbon monoxide.

In addition to that, various isotherms were obtained by Brunner et al. and the data points

of temperatures below 373.15 K and pressures up to 10 MPa were plotted in Figure 2-12.

It shows that with increasing temperature, the solubility of CO in methanol increases

correspondingly although the increments are not significant. This phenomenon is the

same as hydrogen in methanol while different to most gases.

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0.000 0.005 0.010 0.015 0.020 0.025 0.030 0.035 0.0400

1

2

3

4

5

6

7

8

9

10

Pre

ssure

(M

pa)

mole fraction

298.15 K

323.15 K

373.15 K

Figure 2-12. Solubility data of CO-methanol binary system at various temperatures [106]. The x-axis is the mole fraction of H2. The y-axis is the total pressure.

2.3.4 SOLUBILITY MODELLING

Thermodynamic modelling studies in the literature are mostly focused on the H2-

methanol system. Those studies can be summarised as follows:

BINARY INTERACTION PARAMETER (BIP)

Peng and Robinson developed a two-parameter (cohesion parameter a and covolume b)

cubic equation of state (PR EoS). The equation combines simplicity and accuracy and

showed equal or better results than those of Soave-Redlich-Kwong (SRK) equation in all

tested cases and with a major advantage in predicting the liquid-phase densities.

The Peng Robinson Equation of State with a binary interaction parameter kij is usually

fitted to the solubility isotherms to ensure that the Equation of State (EoS) describes the

experimental results correctly, especially for work relying on the static-analytic methods

[109]. The generalised Peng Robinson Equation of States (PR-EOS) is presented in

Equation 2-14.

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Equation 2-14. Generalised equation of Peng Robinson EoS

P =RT

Vm − b−

𝑎

𝑉𝑚(𝑉𝑚 + 𝑏) + 𝑏(𝑉𝑚 − 𝑏)

Where, Vm is the molar volume, a and b are parameters

For a binary system, a and b can be determined by the following equations (Equation 2-15

to Equation 2-18)

Equation 2-15. The parameter b determination of binary system

b = ∑ xibi

NC

i=1

Equation 2-16. The parameter a determination of binary system

a = ∑ ∑ xixj√aiaj(1 − kij)

NC

j=1

NC

i=1

, kii = 0, kij = kji

Equation 2-17. b of the pure component

bi = 0.07780 RTC,i

PC,i

Equation 2-18. a of the pure component

ai = αi × 0.45724 R2TC,i

2

PC,i

Where, x is the mole fraction, and subscript i and c indicate species and critical properties

of the components, respectively. α (alpha function) was introduced to improve vapour

pressure prediction especially for polar fluids. Soave, Twu and Boston – Mathias alpha

function are the three common alfa functions used in industrial applications and studied

in the literature, even though around 20 alpha functions were suggested for various

temperature and pressure ranges[112] [113]. Such alpha functions can be used with PR-

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EoS or SRK-EoS. Descamps et al. applied the Twu alpha function and Gemo et al. employed

the Boston-Mathias alpha function in their studies [108], [109].

(a) Soave alpha function

The alpha function (Equation 2-19) was first introduced by Soave in the Redlich-Kwong

EoS, which improves the correlation of vapour pressure of the pure component [114].

Later on, Peng Robinson suggested an empirical correlation for the alpha function which

is in terms of parameter mi (Equation 2-20) and reduced temperature, where the mi is a

generalised function of acentric factor [115]. Such a function can be applied at any

temperature conditions.

Equation 2-19. Alpha function by Soave

αi(T) = [1 + mi(1 − √Tr,i)]2

Equation 2-20. Generalised function m of acentric functor

mi = 0.37464 + 1.54226ωi − 0.26992ωi2

where, 𝜔 is the acentric factor.

(b) Twu alpha function

Twu alpha function (Equation 2-21) has three parameters which are stated as L,M,N in

their original publication [116]. The values of L, M and N are component-dependence, and

can be determined from regression of pure component vapour pressure.

Equation 2-21. Alpha function by Twu

α = TrN(M−1)

expL(1−TrNM)

In 1995, a generalized Twu equation (Equation 2-22) was proposed with respect to

acentric factor and two alpha functions denoted by 𝛼(0) and 𝛼(1) [117]. For each 𝛼(0) and

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𝛼(1) function, generalized sets of L, M and N are employed which are applicable for all

compounds at subcritical and supercritical conditions which are summarised in Table 2-6.

Equation 2-22. Generalised alpha function by Twu

α = 𝛼(0) + 𝜔(𝛼(1) − 𝛼(0))

Table 2-6. The generalized Twu alpha function parameters for subcritical and supercritical conditions

Generalized Twu Alpha parameters

Tr ≤ 1 Tr > 1

αsub(0)

αsub(1)

αsup(0)

αsup(1)

L 0.141599 0.500315 0.441411 0.032580 M 0.919422 0.799457 6.500018 1.289098 N 2.496441 3.291790 0.200000 8.000000

(c) Boston-Mathias alpha function

In 1980, Boston and Mathias proposed a modification in the original Soave alpha function

for temperatures above TC (Equation 2-23 ) [118]. In this model, alpha function remains

the same as Soave’s when temperatures are below the critical temperature.

Equation 2-23. Alpha function by Boston-Mathias at subcritical condition

Tri ≤ 1, 𝛼𝑖(𝑇) = [1 + 𝑚𝑖(1 − √𝑇𝑟,𝑖)]2

Equation 2-24. Alpha function by Boston-Mathias at supercritical condition

Tri ≥ 1, αi(T) = exp[c(1 − Trid)] , d = 1 +

mi

2, c =

mi

d

mi = 0.37464 + 1.54226ωi − 0.26992ωi2

Therefore, by solving all the equations performed above, the modelling data can be

regressed with the experimental data based on adjusting the binary interaction

parameter only (BIP) [109]. Based on the different experimental systems, this value can

vary to a large extent. For a typical polar solvent, the value of 𝑘𝑖𝑗 is negative or greater

than 0.2, whereas for a non-polar solvent, the value is located between 0 and 0.2 [119].

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In addition, the binary interaction parameter (BIP) kij is dependent on temperature only

and can be expressed by the following equation.

Equation 2-25. Binary interaction parameter (BIP)

kij = kij1 + kij

2 ∙ T

There are two modelling approaches describing the equilibrium behaviour of vapour-

liquid systems: Equation of State (EoS) and activity model. The EoS method takes

advantage of the equation of state for both gas and liquid phases, whereas the activity

model utilises the activity coefficients to describe the liquid phase and uses the equation

of state for the vapour phase. The EoS method has been successfully applied to non-polar

and slightly polar solvents under a broad range of temperatures and pressures [120]. The

activity model is commonly used for incompressible solvents at low temperatures and

pressures [121]. As the reaction conditions of methanol synthesis processes are at

relatively high temperatures and high pressures, the activity model may not be applicable

to the system. In addition, the activity model incorporates a number of interaction

parameters (usually more than 4 parameters), which requires great amount of

experimental data to regress, and thus it may not be accurate due to the lack of available

data for vapour-liquid equilibrium of methanol synthesis processes. Therefore, EoS

model is usually used to predict the phase behaviour of the methanol synthesis processes.

HENRY’S CONSTANT

Thi et al. tried to fit the Henry’s constants to all reliable H2-methanol literature sources in

order to evaluate the Henry’s constant of hydrogen in solvents at high temperature [122].

Two empirical models suggested by Harvey were applied to determine the parameters

for the regressions, and both equations and the corresponding parameters are shown in

Table 2-7, along with the absolute average deviation (AADR). H𝑖 is the Henry’s constant

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of the system, 𝑃𝑆𝜎 is the vapour pressure of the saturated solution, and 𝑇𝑟 is the reduced

temperature. The absolute average deviation is evaluated by the difference between

experimental and calculated values of Henry’s constant, and the equation of AADR is

shown in Equation 2-26. In Table 2-7, the three-parameter equation provides a slightly

better regression compared to the two-parameter equation with about 10% uncertainty.

In addition to that, molecular simulations were also conducted in Thi’s work, the

regression however is not desirable for H2-methanol system, for which AADR gives 84%.

Equation 2-26. Absolute average deviation relatives (AADR)

AADR (%) =1

Npts∑|

𝐻𝑖𝑐𝑎𝑙 − 𝐻𝑖

𝑒𝑥𝑝

𝐻𝑖𝑒𝑥𝑝 |100

Table 2-7. Coefficients of two parameter and three parameter equations in H2-methanol system

Equation Parameters AADR a or a’ b or b’ c

Tr ln (𝐻𝑖

𝑃𝑆𝜎) = 𝑎 + 𝑏(1 − 𝑇𝑟)

2.9188 7.5392 N.A. 9.57%

Tr ln (𝐻𝑖

𝑃𝑆𝜎) = 𝑎′ + 𝑏′(1 − 𝑇𝑟)0.355

+ 𝒄𝑇𝑟 (1

𝑇𝑟− 1)

1.5

0.7503 6.2275 2.1136 9.2%

PSEUDO-HENRY’S CONSTANT

A ‘pseudo-Henry’s law constant, H2,1PS’ introduced by Breman et al. is used to show the

property of the solubility of gases in liquids [105], [123]. Based on the definition of

fugacity, when equilibrium of a gas and liquid are attained, the following equation

established (as discussed earlier).

Equation 2-27. Fugacity equilibrium

fiV = fi

L

The gas phase fugacity can be determined from Equation 2-28.

Literature review

56

Equation 2-28. Vapour phase fugacity of component

fiV = ϕi

VyiP

Liquid phase fugacity at constant temperature and composition depends slightly on

pressure as shown in the following equation. Subscript 1 and 2 stands for solute and

solvent, respectively. The exponential term is called the Poynting correction.

Equation 2-29. Liquid phase fugacity

fiL = fi

L(P2sat) ∙ exp (∫

v1̅

RT

P

P2sat

dP)

For a dilute solution of solute 1 in solvent 2, the liquid phase fugacity of the solute is

usually given by Henry’s law:

Equation 2-30. Liquid phase fugacity in terms of Henry’s law constant

fiL = γ1x1H12 with lim

x1→0(γ1) = 1

Therefore, combining Equation 2-27 to Equation 2-30 we get,

Equation 2-31. The relationship between gas phase fugacity and liquid phase activity coefficient

ϕ1𝑉𝑦1𝑃 = 𝛾1𝑥1𝐻12exp (∫

𝑣1̅̅ ̅

𝑅𝑇

𝑃

𝑃2𝑠𝑎𝑡

𝑑𝑃)

In addition, pseudo Henry’s law constant H12𝑃𝑆 can be written as Equation 2-32.

Equation 2-32. Pseudo Henry’s law constant

H12𝑃𝑆 = 𝛾1𝐻12

Hence, rearranging Equation 2-31 to obtain the formula of H12PS to determine the pseudo

Henry’s law constant by experimental data.

Equation 2-33. The formula of pseudo Henry’s law constant

H12𝑃𝑆(𝑃2

𝑠𝑎𝑡) =ϕ1

𝑉𝑦1𝑃

𝑥1exp (∫

𝑣1̅̅ ̅

𝑅𝑇

𝑃2𝑠𝑎𝑡

𝑃

𝑑𝑃)

Literature review

57

Applying the Peng-Robinson EoS to determine ϕ1𝑉 and v1̅ , and using either Antoine

Equation or Wagner Equation to obtain vapour pressure of the solvent at specific

temperature [124]. In addition, the composition of solute and solvent in each phase can

be calculated from the experimental data at equilibrium using Breman’s iterative method.

Hence, the pseudo-Henry’s law constant can be determined [123].

2.4 CONCLUSIONS

In this chapter, three types of recent methanol synthesis methods have been discussed in

detail including kinetics mechanism, kinetics and roles of catalysts. Solubility

experimental data and thermodynamic regressed models were extensively studied and

investigated for the current research topic in the literature.

As can be seen, due to the insufficient reliable VLE data from previous work, performing

VLE experiments became important and necessary in our study – both to confirm if mass

transfer control was occurring and to obtain values of liquid phase concentrations of CO

and H2 for kinetics modelling. Hence, solubility experiments were conducted for four

systems: H2-methanol, H2-methyl formate, CO-methanol and CO-methyl formate at

temperature ranges from 25 to 100 °C and pressures from 0.3 to 3 MPa. Such operating

conditions were selected based on the reaction conditions. The solubility experimental

apparatus, procedures and the experimental data will be presented in chapter 4.

Materials and methodology

58

CHAPTER 3 MATERIALS AND METHODOLOGY

3.1 MATERIALS

All the chemicals and gases used in this work and their specifications and sources are

listed in Table 3-1 and Table 3-2, respectively. Gases and chemicals were employed as

used without further purification unless otherwise stated.

Table 3-1. Information of chemicals used in the study

Chemicals Formula Suppliers Purity Applications Methanol (Anhydrous) CH3OH Sigma Aldrich 99.9% Chapter 4 and 5 Methyl formate (Anhydrous)

CH3OCOH Sigma Aldrich 99.0 % Chapter 4 to 7

Copper chromite 2CuO Cr2O3 Sigma Aldrich N.A. Chapter 5 Potassium methoxide solution

CH3OK Sigma Aldrich 25.0% in methanol

Chapter 5

Copper nitrate hemi (pentahydrate)

Cu(NO)3·2.5H2O Sigma Aldrich 98% Chapter 6 and 7

Heptane C6H14 Thermo Fisher 99.9% Chapter 5 to 7 Zinc nitrate hexahydrate

Zn(NO)3·6H2O Sigma Aldrich 98.0% Chapter 6 and 7

Zirconium (IV) oxynitrate hydrate

ZrO(NO3)2·3.76H2O Sigma Aldrich 99.0% Chapter 6 and 7

Potassium carbonate (Anhydrous)

K2CO3 Ajex 99.0% Chapter 6 and 7

Hydrotalcite-PURAL® MG50 (MgO:Al2O3(50:50))

Mg2xAl2(OH)4X+4CO3

Sasol Germany GmBH

N.A. Chapter 6 and 7

Table 3-2. Information of gas cylinders used in the study

Cylinders Suppliers Purity Applications Argon (Ar) Coregas Pty.Ltd 99.999% GC analysis Carbon dioxide (CO2) Coregas Pty Ltd 99% Chapter 4 Carbon monoxide (CO) Coregas Pty Ltd 99.995% Chapter 4 and 5 Helium (He) Coregas Pty Ltd 99.999% Leak test Hydrogen (H2) Coregas Pty Ltd 99.999% Chapter 4 to 7 Methane (CH4) Coregas Pty Ltd 99.95% GC calibration Carbon dioxide in helium Coregas Pty Ltd 4.99 vol% CO2 in He TPD-CO2 analysis Hydrogen in argon Coregas Pty Ltd 5.4 vol% H2 in Ar TPR analysis Dimethyl ether in argon Coregas Pty Ltd 4.96% (CH3)2O in Ar GC calibration Instrument air Coregas Pty Ltd 21% O2 in N2 GC calibration Carbon dioxide in helium Coregas Pty Ltd 14.96% CO2 in N2 GC calibration


59

3.2 METHODOLOGIES

3.2.1 POWDER X-RAY DIFFRACTION (XRD)

XRD is an analytical technique used to identify the structure and phase composition of

crystalline catalysts [125], [126]. The diffraction patterns of the synthesised catalysts

were recorded by a Bruker D2 PHASER X-ray powder diffractometer using a nickel

filtered Cu K𝛼 radiation (𝜆=1.5406 Å). The X-ray diffraction scanning was performed at

ambient conditions at 2θ from 10° to 90° at 30 kV and 10 mA using a scan rate of 2°/min.

3.2.2 SCANNING ELECTRON MICROSCOPY (SEM)

The morphology and the size of the produced solid catalysts were determined using a

field emission scanning electron microscopy taken by JSM-7001F Schottky at 5kV [127].

The samples were crushed and sprinkled on carbon tape and mounted on a metal stub

and coated with gold.

3.2.3 ENERGY DISPERSIVE X-RAY SPECTROSCOPY (EDX)

The elemental (both metals and non-metals) analysis as well as their relative proportions

(atomic % generally) were determined using the energy dispersive x-ray spectrometer

attached to the JSM-7001F Schottky [128].

3.2.4 N2 ADSORPTION-DESORPTION ISOTHERMS

The solid catalysts, especially metal oxides, supported metal oxides, are porous materials

with complex pore size distributions that range between micro (average pore diameter d

< 2 nm), meso (2 < d < 50 nm) and macro (d > 50 nm) [129]. The surface area, pore volume,

and pore size distribution play a vital role in the number of accessible active sites in the

catalyst. The structural properties of the solid catalysts were characterised by N2


60

adsorption-desorption measurement at 77K using ASAP 2010 (Accelerated Surface Area

and Porosimetry 2010) [125]. All samples were degassed at 573 K for 3 hours under

vacuum before the analysis. The Brunauer, Emmer, and Teller (BET) method was used to

calculate the specific surface area (SA) using the adsorption data collected over a relative

pressure (P/P0) range of 0.05 – 0.3. The pore volume (PV) was calculated by Barrett-

Joyner-Halenda (BJH) method using the adsorption data.

3.2.5 TEMPERATURE-PROGRAMMED REDUCTION (TPR)

Temperature-programmed reduction (TPR) was used to characterise the metal oxides,

mixed metal oxides, and metal oxides dispersed on a support [130]. In this work, the

oxide of interest is copper oxide. The TPR provides quantitative information of

reducibility of the oxide’s surfaces, as well as the heterogeneity of the reducible surface.

BELCAT Basic, the chemisorption analyser, was used for TPR measurements of the

catalysts samples. Around 40 mg sample was placed on top of glass wool in a U-tube

quartz reactor. The sample was pre-treated by flowing pure Ar (50 mL/min) at 573 K for

2 hours to remove H2O content. Then the sample was cooled down to 323 K by flowing

pure Ar at 50 mL/min. The TPR experiments carried out under 5 vol% H2/Ar flowing at

50 mL/min with temperature increase at a ramp rate of 5 K/min until reach 1023 K.

A mass spectrometry MS (BELMass) was used to measure the change of H2 concentration

in the outlet gas stream over the temperature variation. The BELMass is a quadruple MS

with a faraday cup detector. It operates at 1 mA electron current and 1000 V secondary

electron multiplier. The hydrogen gas (H2) was calibrated prior the TPR experiment in

order to accurately calculate the amount of H2 consumed. The calibration of the hydrogen

peaks with correction factor can be found in Appendix A. CO2 - Temperature programmed

desorption (CO2-TPD)


61

Temperature programmed desorption of carbon dioxide (CO2-TPD) is a method to

determine the catalyst basicity as well as the distribution of the basic sites on the surface

of the catalyst [131]. Pure CO2 gas was supplied for CO2-TPD measurement in the BELCAT

Basic chemisorption analyser. Mass Spectrometry (BELMass) was used to determine the

amount of the actual desorbed CO2. Four procedures are involved in the CO2-TPD

experiment.

Step 1: Catalyst reduction

100 mg of the sample was placed on the top of glass wool in a U-tube quartz reactor where

pure H2 (30 mL/min) flowed at 573 K for six hours and then cooled down to room

temperature using He (30 mL/min).

Step 2: Pre-treatment

After the reduction step, the sample weight was measured to obtain the actual mass. The

sample was pre-treated by flowing pure He (50 mL/min) at 573 K for 2 hours to remove

water content. Subsequently, the sample was cooled down to 393 K by flowing He at 50

mL/min.

Step 3: CO2 adsorption

To determine the amount of the basic sites of the catalysts, only chemisorption is

considered. 393 K was selected as the exposure temperature in the CO2-TCD experiment.

The pre-treated sample was exposed to CO2/He stream (30 mL/min) for 60 mins at 393

K.

Step 4: Chemisorption measurement

Chemisorption is a type of adsorption where the adsorbate molecules interact with the

catalyst surface by strong chemical bonds [125]. The samples were heated from 393 K to

1073 K at a ramp rate of 10 K/min by He stream. Then the samples were held at 1073 K


62

for 2 hours to maximise the residual CO2 desorption. Mass Spectrometry (MS) was used

to detect the CO2 in order to determine its desorption amount. The total amount of the

basic sites can be calculated using Equation 3-1.

Equation 3-1. Determination of amount of desorbed CO2

𝑎𝑚𝑜𝑢𝑛𝑡 𝑜𝑓 𝑑𝑒𝑠𝑜𝑟𝑏𝑒𝑑 𝐶𝑂2 = ∑ 𝑀𝑆𝑠𝑖𝑔𝑛𝑎𝑙𝑠 × 𝐶𝐹

The calibration factor (CF) was determined by performing calibration experiments on the

MS. Up to 20 repeat pulse experiments were conducted and each time identical dosing

volume of 4.99 vol % CO2/He gas was injected at room temperature and atmospheric

pressure. 20 peaks were obtained, and each peak area was calculated with the unit of

intensity*sec. Only the last three peaks were used to calculate the average peak area.

Based on the recorded pressure and temperature of the oven as well as the dosing volume

of the gas mixture, the moles of CO2 was determined based on the ideal gas law. The

correction factor (CF) was computed by dividing the average peak area by the dosing

moles of the gases mixture. The calibration data of CO2 using MS are summarised in

Appendix B.

3.2.6 SPECIFIC COPPER SURFACE AREA VIA N2O TITRATION

The specific surface area of metallic copper and copper dispersion of the copper-based

catalysts were measured by N2O titration. Three sequential steps were involved in the

N2O chemisorption process.

Reaction 3-1. Reduction of copper (II) oxide to metallic copper

CuO + H2 = Cu + H2O and the hydrogen consumption = A1

Reaction 3-2. Oxidation of copper by nitrous oxide to copper (I) oxide

2Cu + N2O = Cu2O + N2


63

Reaction 3-3. Reduction of copper (I) oxide to metallic copper

Cu2O + H2 = 2Cu + H2O and the hydrogen consumption = A2

BELCAT Basic chemisorption analyser was used to conduct the N2O titration experiment.

Step 1 (Reaction 3-1) represents the reduction of CuO phase in the catalyst to yield the

metallic copper. In this step, the sample was placed on the top of the glass wool in the U-

tube quartz reactor and reduced with a flow of 5.4 vol% H2/Ar (30 mL/min) at 573 K for

six hours and then cooled down to room temperature using He (30 mL/min). Step 2 gives

the oxidation of metallic Cu to Cu2O by N2O, which is the main step in the N2O titration, to

determine the metallic copper dispersion. N2O (30 mL/min) was introduced to the

catalyst at 333 K for 30 minutes and then the catalysts were purged with pure He for 30

minutes to remove the residual N2O. Step 3 is the reduction of the Cu2O to metallic Cu

with 5.4 vol% H2/Ar (30 mL/min). The temperature was increased to 673 K with a

heating rate of 10 K/min. The dispersion of Cu was defined by Equation 3-2 and the

specific area of metallic copper was determined from H2 consumption (A2) with 1.46 ×

1019 copper atoms per m2.

Equation 3-2. Determination of copper dispersion

D =2A2

A1× 100%

3.2.7 THERMAL GRAVIMETRIC ANALYSIS (TGA)

Netzsch TG 209 F1 Libra was used to measure thermal gravimetry of the samples and

integrated FTIR was able to provide the information of the evolute gases.

The dried sample (around 12 mg) was placed in the alumina crucibles and the thermal

analysis was performed in pure nitrogen at a heating rate of 2 K/min in the range of 303

K and 1073 K.


64

3.2.8 X-RAY PHOTOELECTRON SPECTROSCOPY (XPS)

X-ray photoelectron spectroscopy (XPS) is a technique for elemental composition

analysis of the surface of the solid catalysts and for the oxidation state and electronic of

the predetermined elements [132].

A Kratos Axis ULTRA X-ray Photoelectron Spectrometer equipped with a 165mm

hemispherical electron energy analyser was used to acquire XPS data. The incident

radiation was Monochromatic Al KαX-rays (1486.6 eV) at 150 W (15 kV, 15 mA). The base

pressure in the analysis chamber was set at 1.0 ×10-9 torr and the pressure of the sample

analysis was maintained at 1.0 ×10-8 torr. The scanned area was approximately 0.8 mm ×

0.3 mm and the depth was less than 10 nm (the volume is c.a. 2400 µm3). The survey

(wide) scans were taken at an analyser pass energy of 160 eV and were carried out over

the range of 1200 to 0 eV binding energy range with 1.0 eV increment and a dwell time of

100 ms. The multiplex (narrow) high resolution scans at a pass energy of 20 eV and were

run with 0.05 eV step and 250 ms dwell time.

3.2.9 AUGER ELECTRON SPECTROSCOPY (AES)

Auger electron spectroscopy (AES) is obtained from the ejection of the Auger electron

after relaxation of the photoionized atoms [132]. This technique is complementary to the

XPS results and provides additional surface-sensitive information on the surface

compositions and the specific chemical bonding. The operating conditions are same as

that of XPS.

3.2.10 PRODUCTS ANALYSIS METHOD - GAS CHROMATOGRAPHY (GC)

Gas chromatography (GC) is an instrumental technique that is used to analyze the

compositions of mixtures by separating the mixtures into individual components.


65

In our present research, both thermal conductivity detector (TCD) and flame ionization

detector (FID) were employed. A configuration of GC, Agilent 7890B, was designed based

on the feed composition and the expected reaction mixture components, including H2, CO,

CO2, CH4, O2, N2, Ar, (CH3)2O, alcohols (C1-C4), CH3OCOH, and HCOOH. Two columns in

series were installed in the TCD line, where the column CP-Molesieve 5 Å was used to

separate CO2, and the other column HP-PLOT-U column was used for the separation of H2,

CO, air, CH4, (CH3)2O and organic compounds. In addition, a column of HP-INNOWax was

installed in the FID line to separate organic products. An automatic liquid sampler was

designed to inject the liquid samples (1 µL) into the vaporiser. The gases in the sampler

tank was injected to the GC manually via a needle valve attached to the GC inlet. The

sampled gases ware initially purged to the GC for 1 minute to remove the gases residues

in the GC lines. A volume of 250 µL of each gases sample was then injected into the column.

Each analysis took 22.5 minutes. The schematic diagram of the valve and detectors in the

GC is shown in Figure 3-1. Since the liquid samples were mainly methanol and methyl

formate, the calibration was carried out based on multi-point calibration by injecting a

series of liquid methanol and methyl formate mixtures with known different

compositions. The calibration calculation of these two compounds can be found in

Appendix C. A ScottTM analytical bottle with number of premixed gases components was

used to identify the retention time of the gases components (The details of the

composition of gases mixtures are shown in Appendix D). The retention time of gases

including CO and (CH3)2O were determined using pure CO and 4.96% (CH3)2O in Ar

mixture, respectively. Calibration of the gases sample was accomplished by a single-point

calibration method and the details of the calibration lines can be found in Appendix D.

• Pure gases: N2, CO, H2, CH4, and CO2

• Mixture gases: Compressed air and 4.96% (CH3)2O in Ar


66

Figure 3-1. Schematic diagram of the valves and detectors in the Agilent GC 7890B

Solubility study

67

CHAPTER 4 SOLUBILITY STUDY

4.1 OBJECTIVE

In the literature review (Chapter 2), a two-step methanol synthesis via methyl formate

was proposed for investigation in the current research study. Prior to investigating the

kinetics of the two step reactions, it is of great importance to understand the solubilities

of the reactant gases (carbon monoxide and hydrogen) into liquid solvents (methanol and

methyl formate) to evaluate the effect of the mass transfer in the reaction process if the

reaction is controlled by the mass transport process and to understand the equilibrium

limits, if any, of the gases in the liquids.

In this chapter, a validation experiment of CO2 solubility in methanol was carried out to

verify the apparatus and procedure by comparing the results with the literature data. In

addition, the solubilities of CO and H2 in liquid methyl formate and methanol were

measured at different pressures (0.3 MPa to 3.3 MPa) and temperatures (25 °C to 100 °C),

respectively. The Peng-Robinson Equation of State (PR-EoS) was used to fit and validate

the experimental data.

4.2 EXPERIMENTAL APPARATUS AND PROCEDURES

4.2.1 APPARATUS

The apparatus used for solubility measurement in this work was designed based on the

constant-volume methodology, where a measured amount of gas (solute) was brought

into the vessel and contacted with a known volume of pure liquid (solvent) via stirring

until the equilibrium was achieved. The solute solubility in the solvent can be calculated

from fundamental thermodynamic equations.

Solubility study

68

The solubility measurement apparatus is shown in Figure 4-1. It consists of a gas cylinder

1, mass flow controller 2 (Brooks® Instrument 5850E series), a storage tank 3

(Swagelok® Stainless steel sample cylinder), an absorption tank 4 (Swagelok® Stainless

steel sample cylinder), a heating tape 5 (BriskHeat®), a magnetic stirrer 6 (Industrial

Equipment & Control Pty. Ltd), and a vacuum pump 7 (Elnor Motors®). The volume of the

storage tank 3 and the absorption tank 4 are 183.3 cm3 and 82.5cm3, respectively,

including connecting tubes.

Figure 4-1. Schematic diagram of solubility apparatus: 1. Gas cylinders (He, CO2, H2 and CO); 2. Mass flow controller, 3. Storage tank; 4. Absorption tank; 5. Heating tape/Cooling bath; 6. Magnetic stirrer; 7. Vacuum pump; 8. Vent system; BV-1 to BV-4: Ball valves; NV-1 and NV-2: Needle valves

Investigated gases are charged from the gas cylinder 1 into the storage tank 3 by

manipulating the mass flow controller 2. The storage vessel serves the purpose of the

determination of the amount of gases fed into the absorption tank via evaluating the

pressure difference of the storage tank before and after loading the gas.

Pressures of the storage tank and the absorption tank are monitored using a Swagelok S-

model pressure transducer (accuracy ≤ 0.25% span limit point calibration) and a MKS

Type 627B pressure transducer (accuracy ± 1 mmHg), respectively. Both pressure

TT A

7

1

6

BV-1 BV-2 BV-3 BV-4

NV-1 NV-2

2

3

PT A PT B TT B

4

5

8 8

Computer

Solubility study

69

transducers are connected to a NI™ USB-6002 data acquisition (DAQ) device and

monitored by the LabVIEW System Design Software. Meanwhile, the pressure of the

absorption tank B is displayed by a power supply digital readout unit (MKS Instruments

Inc, 660B model). Two pressure transducers are calibrated using accurate manometers

to provide precise and reproducible results at a constant operating temperature of 294.6

K.

Temperature inside the absorption tank is controlled using the external heating tape 5

coupled with a temperature controller. Temperatures of both storage tank and

absorption tank are measured using the K-type thermocouples (accuracy ± 0.1 K)

connected to the National Instruments™ (NI™) thermocouple measurement devices and

recorded with the LabVIEW System Design Software.

4.2.2 PROCEDURE

The procedure of vapour-liquid equilibrium measurement at moderate temperatures and

moderate pressures includes the following steps:

1) Leak Test: Prior to every experiment, a leak test is undertaken by injecting Helium to

a certain pressure value and monitoring the stability of the pressure value over two hours.

2) Liquid solvent load and system degas: 12 mL liquid solvent (methanol or methyl

formate) is loaded to the absorption tank 4 under a pure nitrogen environment in a glove

box. The tank is subsequently connected to the process apparatus line via a VCR-gasket

(Swagelok®), and then immersed in dry ice to cool down. The purpose of using dry ice is

to reduce the vapour pressure to an acceptable value (less than 500 Pa) due to the

volatility of solvents. Once the temperature inside of the absorption tank is reduced

below 213 K and the vapor pressure of the volatile solvents is below 200 Pa, a vacuum

Solubility study

70

pump is turned on to degas the apparatus. The solvent loss during the degas step is less

than 0.5 % based on the mass balance calculation, and it is assumed that there is no other

gas contained in the absorption tank after degassing.

3) Gas injection: While waiting for the solvent in the absorption tank 4 to cool down to

213 K, valves BV-1 and BV-2 are opened to load the gas phase solute from gas cylinder to

the storage tank 3 via the mass flow controller (MFC).

4) Solubility experiment: when the pressure inside the storage tank 3 is stabilized, the

valve BV-3 is opened for 3 seconds and then closed, which allows the gas to flow from the

tank 3 to the tank 4. Subsequently, the dry ice is removed, and we wait for the tank 4 to

reach room temperature. Then the heating tape (No. 5 in Figure 4-1) is attached to the

tank 4 and heated up to the preset temperatures, simultaneously the stirrer is turned on

to accelerate the gas-liquid mass transfer. When the pressure measured by the pressure

gauge transducer (PT B in Figure 4-1) is stable, the system is considered to reach the

equilibrium state. In general, the system takes one hour to reach equilibrium.

After the stabilized pressure has been recorded, valves BV-1 and BV-2 are reopened to

load more gas phase solute into the storage tank 3, and then closed again. Thereafter,

valve BV-3 is opened for only 3 seconds to pressurize the tank 4. A new equilibrium

pressure is then achieved after one hour. The steps mentioned in this paragraph are

repeated to get a number of equilibrium pressure values in the tank 4. The final recorded

value is about 3.2 MPa. All pressure and temperature values of the tank 3 and the tank 4

are continuously recorded every 1 second via LabVIEW software.

5) Shut down procedure: When the experiment is completed, the tank 4 is cooled down

to 213 K using dry ice to reduce the vapor pressure inside the tank. A venting step is then

Solubility study

71

conducted by opening the valve NV-1 and NV-2, followed by detaching the tank 4. The

volume of the solvent in tank 4 is measured. A flow chart of the experimental procedure

is shown in Figure 4-2. The mole fraction of the gas phase solutes in the solvents can be

calculated from the material balance of components and the equations of state.

Experiments are carried out at 298.15 K, 323.15 K, 348.15 K and 373.15 K with a variation

of ± 0.1 K.

Figure 4-2. The flow chart of solubility experiments

Tank 3

Check leakage

Introduce gas solute from gas cylinder Liquid solvent injection

Stabilize P & T

Cool down using dry ice

Load gas from tank 3 to tank 4

Tank 4

Degas by vacuum pump

Stabilize P & T

Introduce gas solute from gas cylinder

Stabilize P & T Equilibrium estabilish

Heat up to pre-set T

Release solute to vent system Cool down using dry ice

Release solute to vent system

Measure volume of solvent

Shu

t-d

ow

n S

tep

Exp

erim

ent

rep

eati

ng

step

sSt

art-

up

Ste

p

Rep

eat

step

s Rep

eat steps

Solubility study

72

4.3 THEORY

4.3.1 EVALUATION OF EXPERIMENTAL RESULTS

The total amount of the solute fed into the absorption tank 4 is calculated from Equation

4-1. The values of n1 and n2 can be obtained from the Peng-Robinson cubic equation of

state (PR EoS) [133].

Equation 4-1. The total amount of solute in the tank 4

nsoluteT = n1 − n2

where, n1 and n2 are the amount of gas solute in the storage tank 3 before and after the gas is fed into tank 4.

The PR equation is given in Equation 4-2. Since the gas in the storage tank 3 is pure,

parameters a and b can be determined based on the critical conditions of pure

components from Equation 4-3 and Equation 4-4 which are tabulated in Table 4-1[134].

α is a function (given in Equation 4-5) to improve vapour pressure prediction, especially

for polar fluids [117]. The parameter m in Equation 4-5 can be determined from Equation

4-6 to Equation 4-7 [135]. ω is the acentric factor of the components, which are also given

in Table 4-1.

Equation 4-2. Generalised equation of Peng-Robinson EoS

p =RT

Vm − b−

a

Vm(Vm + b) + b(Vm − b)

where, Vm is the molar volume, a and b are parameters

Equation 4-3. Parameter b in PR EoS

b = 0.07780 RTC

PC

Equation 4-4. Parameter a in PR EoS

a = α ∙ 0.45724 R2TC

2

PC

Solubility study

73

Equation 4-5. alpha function

α(T) = [1 + m(1 − √Tr)]2

Equation 4-6. Determination of mi when acentric factor less than 0.49

m = 0.37464 + 1.54226ω − 0.26992ω2 for ω < 0.49

Equation 4-7. Determination of mi when acentric factor above 0.49

m = 0.376942 + 1.48503ω − 0.1644ω2 + 0.016667ω3 for ω > 0.49

Table 4-1. Physical properties of pure components

Components MW (g/mol) Tc (K) Pc (MPa) Vc (m3/mol) Zc Acentric factor CH3OH 32.04 512.58 8.0959 0.1178 0.224 0.5656 CH3OCOH 60.05 487.20 5.9984 0.172 0.255 0.2537 H2 2.016 33.25 1.297 0.06503 0.305 -0.2153 CO 28.01 132.92 3.4988 0.0931 0.295 0.0663 CO2 44.01 304.21 7.383 0.094 0.274 0.223621

The total amount of solvents fed into the absorption tank 4 can be determined using

Equation 4-8. The density can be calculated using the DIPPR equation (shown in Equation

4-9), where the parameter A, B, C and D are listed in Table 4-2.

Equation 4-8. Total moles of solvents in the equilibrium cell

nsolventT =

ρVL

MW

where, ρ is the density of the solvent at ambient temperature, MW is the molecular weight of the solvent, VL is the liquid volume.

Equation 4-9. The density of the solvent

ρ =A

B1+(1−TC

)D × MW

Table 4-2. The parameters for the solvent density determination

Solvents A B C D Methanol 1.2057 0.19779 512.63 0.17272 Methyl formate 1.1639 0.23213 497.22 0.23826

The total pressure in the absorption tank 4 is the sum of the partial pressure of the gas

phase solute and the partial pressure of the liquid solvent, which can be written in

Solubility study

74

Equation 4-10. The saturated vapour pressure can be determined using the Antoine

equation (given in Equation 4-11), where the parameters are indicated in Table 4-3.

Equation 4-10. The total pressure in the absorption tank 4

PT = Psolute + xsolventPsolventsat

Equation 4-11. The Antoine Equation

log10 Psolventsat = A −

B

T + C − 273.15

Table 4-3. The parameters for the solvent saturated pressure

Solvents A B C Methanol 5.20277 1580.08 239.5 Methyl formate 4.29529 1125.2 230.56

In addition, the total mole of the liquid solvent and the gas phase solute in the tank 4 can

also be written using Equation 4-12 and Equation 4-13. The amount of the solvent in the

vapour phase can be determined using the Peng-Robinson EoS (Equation 4-14). The mole

fraction of the solute in the liquid phase is thereafter calculated using Equation 4-15.

Equation 4-12. The total amount of the solvent in the tank 4

nsolventT = nsolvent

L + nsolventV

Equation 4-13. The total amount of the solute in the tank 4

nsoluteT = nsolute

L + nsoluteV

Equation 4-14. The amount of solvent in the vapour phase

nsolventV = Peng − Robinson (Psolvent, T)

Equation 4-15. The mole fraction of solute in solvent

xsolute =nsolute

L

nsolventL + nsolute

L

4.3.2 MODELLING

Considering this is a binary mixture system, a modified Peng Robinson equation of state

has been used to model the VLE of the H2 and CO system. The generalised equation of the

Solubility study

75

Peng-Robinson Equation of states (PR-EOS) for the mixture is the same with the PR

equation for the pure component (Equation 4-2); however for a binary system, 𝑎 and 𝑏

can be determined using the following equations (from Equation 4-16 to Equation 4-19).

Equation 4-16. Determination of b parameter in PR-EoS

b = ∑ xibi

n

i=1

Equation 4-17. Determination of a parameter in PR-EoS

a = ∑ ∑ xixj√aiaj(1 − kij)

n

j=1

n

i=1

, kii = 0, kij = kji

Equation 4-18. Determination of bi parameter in PR-EoS

bi = 0.07780 RTC,i

PC,i

Equation 4-19. Determination of ai parameter in PR-EoS

ai = αi ∙ 0.45724 R2TC,i

2

PC,i

where, 𝑥 is the mole fraction, and subscript i and c indicate the component and the critical property of the component, respectively.

The 𝑘𝑖𝑗 in Equation 4-17 is an interaction parameter, and the value is fitted based upon

the experimental VLE data. Generally, the 𝑘𝑖𝑗 varies between 0 and 0.2 for nonpolar

solvents and can have a negative or a larger (> 0.2) value for polar species [136]. In the

current research, the binary interaction parameter for H2 and CO system can be

determined using a homogenous approach (phi-phi model), where the fugacity in both

liquid and vapour are equal as indicated in Equation 4-20.

Equation 4-20. Fugacity equilibrium

fiL = fi

V

Solubility study

76

In the literature, the vapour-liquid equilibrium can be described using either a “phi-phi”

approach (using an equation of state to calculate fugacity coefficient for each phase) or

“gamma-phi” approach (using a liquid activity coefficient model for the liquid phase and

the equation of state for the vapour phase). The gamma-phi approach is the traditional

approach which is usually applied at low pressures ranging from ambient pressure to 0.3-

0.5 MPa, whereas the phi-phi approach can be applied at moderate and high pressures

[137]. Since our study mainly focuses on a moderate pressure range of 0.3 MPa to 3 MPa,

the phi-phi approach was adopted. There are six iteration steps involved to evaluate the

binary interaction parameter using the phi-phi approach, and a description flow chart is

given in Figure 4-3 [138].

Given Zi, P, T

Assume the initial Ki

Flash calculation

Calculate KiCalculate

vapour phase Fugacity coefficient

Calculate liquid phase

Fugacity coefficient

Test for convergence

Solution

xi,nLyi,nV

No

Yes

Solubility study

77

Figure 4-3. The flow chart of phi-phi approach to determine the VLE data

Step 1: Determine the initial equilibrium ratio (𝐾𝑖𝐴)

An initial value of the equilibrium ratio (Κ𝑖𝐴) for each component in the mixture at a

specific temperature and pressure is assumed. The initial K value is determined based on

the Wilson’s equation (Equation 4-21).

Equation 4-21. Wilson’s equation

ΚiA =

Pci

Pexp [5.37

1 + ωi

1 −Tci

T

]

where, ΚiA is the assumed initial equilibrium ratio of component i.

Step 2: Determine the mole fraction of the mixture

Perform flash calculations using the assumed Κ𝑖𝐴 values to determine the mole fraction

and the moles of each component in both vapour and liquid phase, 𝑥𝑖 , 𝑦𝑖, 𝑛𝐿 and 𝑛𝑉 . For a

binary system, the composition of the liquid phase can be determined using Equation

4-22 to Equation 4-24.

Equation 4-22. The expression of the mole fraction of components in the liquid phase

∑ xii

= x1 + x2 = 1

Equation 4-23. The expression of the mole fraction of components in the gas phase

∑ yii

= y1 + y2 = K1x1 + K2x2 = 1

Equation 4-24. Determination of the mole fraction of components in terms of equilibrium ratio K

x1 =1 − K2

K1 − K2

Step 3: Determine the fugacity coefficient in the liquid phase

Solubility study

78

The calculated composition of the liquid phase 𝑥𝑖 is used to determine the fugacity

coefficient Φ𝑖𝐿 of each component in the liquid phase. The evaluation expressions are

shown from Equation 4-25 to Equation 4-28.

Equation 4-25. Fugacity coefficient of components in the liquid phase

ln(ΦiL) =

bi(ZL − 1)

bm− ln(ZL − B) − [

A

2√2B] [

2ΨiL

(aα)mL

−bi

bmL

] ln [ZL + (1 + √2)B

ZL − (1 − √2)B]

Equation 4-26. Determination of the mixture parameter Ѱ in the liquid phase

ΨiL = ∑[xj√aiajαiαj(1 − kij)]

j

Equation 4-27. Determination of the mixture parameter 𝑎𝛼 in the liquid phase

(aα)mL = ∑ ∑[xixj

ji

√aiajαiαj(1 − kij)]

Equation 4-28. Determination of the mixture parameter b in the liquid phase

bmL = ∑(xibi)

i

Step 4: Determine fugacity coefficient in the gas phase

Use the calculated composition of the gas phase 𝑥𝑖 to determine the fugacity coefficient

Φ𝑖𝐺 of each component in the gas phase.

Equation 4-29. Fugacity coefficient of components in the gas phase

ln(ΦiV) =

bi(ZV − 1)

bm− ln(ZV − B) − [

A

2√2B] [

2ΨiV

(aα)mV −

bi

bmV ] ln [

ZV + (1 + √2)B

ZV − (1 − √2)B]

Equation 4-30. Determination of the mixture parameter Ѱ in the gas phase

ΨiV = ∑[yj√aiajαiαj(1 − kij)]

j

Equation 4-31. Determination of the mixture parameter 𝑎𝛼 in the gas phase

Solubility study

79

(aα)mV = ∑ ∑[yiyj

ji

√aiajαiαj(1 − kij)]

Equation 4-32. Determination of the mixture parameter b in the gas phase

bmV = ∑(yibi)

i

Step 5: Determine the new equilibrium ratio 𝐾

Determine a new set of equilibrium ratios from Equation 4-33 using the values obtained

from step 3 and step 4.

Equation 4-33. Evaluation of a new equilibrium ratio K

K =Φi

L

ΦiV

Step 6: Test for convergence

Check the solution by applying the constrains that is given in Equation 4-34.

Equation 4-34. Convergence constrains

∑ [Ki

KiA

− 1]

2n

i=1

≤ ε

where, ε is the error tolerance, in this case, we choose 0.00001

n is the number of components in the system

4.3.3 UNCERTAINTY CALCULATION

The uncertainties of the measurement should be constrained within the system errors at

different temperatures, pressures, and volumes. The experimental errors for the

temperature, the pressure and the volume are u(T) =0.2 K, u(Pstorage tank) = 0.001 MPa,

u(Pequilibrium tank) = 200 Pa, and u(V) = 0.02 mL, respectively. Based on the method for the

estimation of uncertainties, the overall uncertainty for the measured solubility can be

determined using Equation 4-35, Equation 4-36 and Equation 4-37.

Solubility study

80

Equation 4-35. Uncertainty of u(x)/x

u(x)

x= √(

u(ng)

ng)

2

+ (u(ng + nl

ng + nl)

2

= √u(n1)2 + u(n2)2 + u(nE)2

ng2

+u(n1)2 + u(n2)2 + u(nE)2 + u(nl)2

(ng + nl)2

Equation 4-36. Uncertainty of u(n)/n in the gas phase

u(n1)

n1=

1

R√(

u(P1)

P1)

2

+ (u(V1)

V)

2

+ (u(T1)

T1)

2

Equation 4-37. Uncertainty of u(n)/n in the liquid phase

u(nl)

n1=

ρ

MWu(VL)

4.3.4 HENRY’S LAW CONSTANT AND ITS CONFIDENCE INTERVALS

The solubility of the gas can be expressed using a form of the Henry’s constant, which is

given in Equation 4-38 [119].

Equation 4-38. The Henry’s constant expression

Hx(P, T) = limxsolute→0

fsoluteliq

(P, T, xsolute)

xsolute

where, Hx (P,T) (MPa) is the Henry’s constant based on the mole fraction, xsolute is the mole

fraction of the gas phase solute in the liquid solvent, fsoluteliq

(P, T, xsolute) is the fugacity of the

solute in the the liquid phase.

At equilibrium, the fugacity of solute in liquid and vapour should be equal, therefore,

Equation 4-39. The fugacity of the solute in the liquid phase

fsoluteliq (P, T, xsolute) = fsolute

vap (P, T, xsolute) = ysolutePϕsolute(P, T, ysolute)

where, fsolutevap (P, T, xsolute) is the fugacity of the solute in the vapour phase, ysoluteis the

mole fraction of solute in the vapour phase, and ϕsolute is the fugacity coefficient of solute in the vapour phase.

Hence, the solute solubility can be determined by Equation 4-40.

Solubility study

81

Equation 4-40. The expression of the Henry’s constant

Hx(P, T) = limxsolute→0

ysolutePϕsolute(P, T, ysolute)

xsolute

To ensure the accuracy of the regressed value, 95% confidence interval can be

determined based on the critical value and standard error.

4.3.5 THERMODYNAMIC PROPERTY DETERMINATION

The thermodynamic properties of the system, including the Gibbs free energy, the

enthalpy of dissolution, and the entropy of dissolution, can be determined using the

experimental and modelling results conducted in this work.

The value of dissolution enthalpy can be defined in Equation 4-41, where ln𝐻𝑥(𝑇, 𝑃) is

expected to have a linear relationship with 1

𝑇.

Equation 4-41. Dissolution enthalpy of gas-liquid solubility

∆disH = R [∂lnHx(T, P)

∂ (1T)

]

P

The dissolution entropy ∆𝑑𝑖𝑠𝑆 can be determined from the intercept of the linear

relationship of ln𝐻𝑥(𝑇, 𝑃) and 1

𝑇, which is given in Equation 4-42. ∆𝑑𝑖𝑠𝑆 and ∆𝑑𝑖𝑠𝐻 are

assumed constant in the investigated temperature range.

Equation 4-42. Dissolution entropy of gas-liquid solubility

∆disS = −R × intercept

Therefore, the dissolution Gibbs free energy ∆𝑑𝑖𝑠𝐺 can be determined using Equation

4-43.

Equation 4-43. Dissolution Gibbs free energy of gas-liquid solubility

∆disG = ∆disH − T ∙ ∆disS

Solubility study

82

4.4 DATA ANALYSIS

4.4.1 VALIDATION OF THE EXPERIMENTAL APPARATUS

Before conducting the solubility experiments, it is necessary to test the reliability of the

experimental apparatus for the current system. Since the vapour-liquid equilibrium of

CO2-methanol system has been extensively studied and a great amount of data is available

in the literature, a validation experiment was carried out using CO2 dissolved into the

liquid methanol in our experimental apparatus to validate the experimental setup and

the data processing method discussed in Section 4.3.

The VLE data for CO2-methanol system was measured at four different temperatures

(300.78 K, 322.91 K, 351.33 K and 376.61 K) with a pressure range of 0.2 MPa to 2.8 MPa.

The results are listed in Table 4-4 and the data are plotted in Figure 4-4.

Table 4-4. Partial pressure (PCO2), liquid phase mole fraction (xi), and uncertainties (δ) of CO2 in methanol from 298.15 K to 373.15 K

𝐏𝐂𝐎𝟐 (MPa) 𝐱𝐢 𝛅 𝐏𝐂𝐎𝟐

(MPa) 𝐱𝐢 𝛅

T=300.78 K T=322.91 K 0.2477 0.0160 0.0004 0.3192 0.0147 0.0004 0.4856 0.0315 0.0007 0.5276 0.0244 0.0007 0.8526 0.0551 0.0006 0.9124 0.0414 0.0006 1.2915 0.0843 0.0008 1.3560 0.0614 0.0008 1.7528 0.1164 0.0010 1.8353 0.0839 0.0010 2.2281 0.1506 0.0012 2.3308 0.1081 0.0010 2.7119 0.1847 0.0014 2.8268 0.1338 0.0014

T=351.33 K T=376.36 K 0.3175 0.0128 0.0004 0.6228 0.0210 0.0019 0.4897 0.0177 0.0007 0.8078 0.0251 0.0001 0.8820 0.0286 0.0006 1.2057 0.0338 0.0008 1.3475 0.0414 0.0009 1.6836 0.0448 0.0011 1.8406 0.0556 0.0012 2.1908 0.0553 0.0014 2.3573 0.0723 0.0015 2.7128 0.0695 0.0017 2.8797 0.0908 0.0017

Solubility study

83

Figure 4-4. The comparison of the experimental results with the literature data

Since the published gas-liquid equilibrium experiments for CO2-methanol system were

performed at various operating conditions, it is difficult to obtain the experimental data

from the literature at the exactly same operating conditions as our work. Hence, to ensure

high accuracy and consistency, the published VLE data that were measured at the

temperatures close to our experimental conditions were selected [104], [106], [139]–

[142]. As can be seen from Figure 4-4, both the measured data in the apparatus and the

literature data of CO2-methanol system have good agreement at these four different

temperatures.

Moreover, as the solubility of CO2 in methanol is larger compared with the solubility of

other gases such as H2 and CO in methanol, it was necessary to test the sensitivity of the

Solubility study

84

apparatus for low solubility systems. Therefore, based on the available CO-methanol VLE

data from Brunner et al. [106], the reliability of using the experimental apparatus to

measure the VLE of CO-methanol system at similar temperatures was checked, and the

results are indicated in Figure 4-5 and Figure 4-6. As can be seen, the solubility results of

CO in methanol obtained from our experimental apparatus are consistent with the

published data, which have the same slope with the literature data. The experimental VLE

results for CO-methanol system conducted in this work are good extensions and

complementary of the existing VLE literature data.

Based upon the study of those comparisons, it can be concluded that the experimental

apparatus to measure the VLE data for CO-methanol, CO-methyl formate, H2-methanol

and H2-methyl formate system in the present study is reliable and of acceptable accuracy.

Figure 4-5. The comparison of the experimental results and literature data for CO-CH3OH system at 298.1 K

Solubility study

85

Figure 4-6. The composition of the experimental results and literature data for CO-CH3OH system at 322.7 K

4.4.2 EXPERIMENTAL RESULTS

Vapour-liquid equilibrium solubility experiments for four systems, including carbon

monoxide in methanol, carbon monoxide in methyl formate, hydrogen in methanol and

hydrogen in methyl formate, were performed at four different temperatures

(approximate 298.15 K, 323.15 K, 348.15 K and 373.15 K).

CO-METHANOL SYSTEM AND CO-METHYL FORMATE SYSTEM

The results of the VLE data for the CO-methanol system and the uncertainties calculated

from Equation 4-35 are summarised in Table 4-5. The Figure 4-7 shows the isotherms of

the phase equilibria of the CO and methanol mixtures. As can be seen, at each specific

temperature, the CO mole fraction in methanol increases with the partial pressure of CO,

but the mole fraction is very small compared with CO2 in methanol at the same partial

pressure, indicating that the solubility of CO in methanol is relatively small, and the

Solubility study

86

Henry’s constant for the CO-methanol system is significantly larger (shown in Table 4-5)

than that of CO2 in methanol. It is noted that the VLE of the CO-methanol system presents

some interesting trends when changing the temperature. The solubility of CO increases

with the increase of the temperature, which is opposite of the typical trend where gas

solubility decreases with increasing the temperature. In most cases, high temperatures

increase the kinetic energy and the Brownian motion of the solute molecules in the liquid

phase, and this increase in thermal energy (kT) exceeds the attractive solute-solvent

interaction, leading to a drop in solubility as temperature increases. However, this

explanation is not applicable for the CO-methanol system. From Figure 4-7, it can be seen

that at the same partial pressure of CO, when increasing the temperature, the liquid mole

fraction of CO in methanol increases, which means that kinetic energy (kT) of the

dissolved CO molecules may not be the dominant factor of the VLE in the CO-methanol

system, and this is also reflected by the Henry’s constants shown in Table 4-5, where the

Henry’s constant decreases with increasing temperature. The phenomenon has also been

found in other published works [102], [106], [123]. One explanation is that CO has

stronger solute-solute interactions than solute-solvent interactions, and they prefer to

associate with themselves rather than the solvents. As the temperature increases, the

thermal expansion of the liquid reduces the opportunities for solute-solvent interactions

and increases the opportunities for solute-solute interactions, leading to an increase in

solubility of CO gas solutes. Although the solvent still expands with increase in

temperature for these systems, this effect can be negligible due to the strong interaction

between the gas molecules and the solvent. As a consequence of these relative energetic

interactions, CO dissolves into methanol and methyl formate are endothermic processes,

where the sum enthalpy of the separation of gas molecules and the separation of liquid

molecules is greater than the enthalpy of mixing gas molecules and liquid molecules. High

Solubility study

87

temperatures provide high energy to break solute bonds in the gas phases, facilitating the

dissolving process of the gas solutes with broken bonds into the solvent.

Table 4-5. Partial pressure (PCO), liquid phase mole fraction (xi), Henry’s law constant (H) and uncertainties (δ) of CO in methanol from 298.15 K to 373.15 K

𝐏𝐂𝐎

(MPa) 𝐱𝐢 𝛅

H (MPa)

𝐏𝐂𝐎

(MPa) 𝐱𝐢 𝛅

H (MPa)

T=298.10 K T=322.71 K

0.2415 0.000848 0.0002

279.2 ±7.42

0.2605 0.000928 0.0001

247.3 ±5.40

0.6093 0.001965 0.0002 0.6005 0.002377 0.0003 1.0820 0.003580 0.0004 1.0737 0.004147 0.0003 1.5870 0.005529 0.0005 1.5876 0.006419 0.0005 2.1011 0.007696 0.0007 2.1099 0.008251 0.0005 2.6167 0.009268 0.0009 2.6420 0.010217 0.0009 3.1489 0.011455 0.0011 3.1672 0.011981 0.0010

T=347.76 K T=372.90 K

0.2878 0.001090 0.0002

219.2 ±5.36

0.5087 0.002368 0.0004

190.49 ±4.28

0.5988 0.002490 0.0003 0.9689 0.004644 0.0003 1.0754 0.004600 0.0004 1.4894 0.007346 0.0005 1.5919 0.007017 0.0005 2.0250 0.010299 0.0009 2.1333 0.009801 0.0011 2.5623 0.013229 0.0013 2.6801 0.012182 0.0012 3.1180 0.015878 0.0011 3.2322 0.014989 0.0015

Figure 4-7. Isothermal phase equilibrium of CO in methyl formate

Solubility study

88

A similar trend has been found in the CO-methyl formate system, and results are shown

in Table 4-6 and Figure 4-8. When increasing the temperature, the solubility of CO in

methyl formate also increases. Compared to the CO solubility in methanol, CO is more

soluble in methyl formate due to the polarity of the solvent. The Henry’s constant also

implies the solubility of CO in methyl formate is greater than that of CO in methanol as

the Henry’s constant for the CO-methyl formate system is smaller than that of CO-

methanol system at the same temperature.

Table 4-6. Partial pressure (PCO), liquid phase mole fraction (xi), Henry’s law constant (H) and uncertainties (δ) of CO in methyl formate from 298.15 K to 373.15 K

𝐏𝐂𝐎

(MPa) 𝐱𝐢 𝛅

H (MPa)

𝐏𝐂𝐎

(MPa) 𝐱𝐢 𝛅

H (MPa)

T=299.72 K T=322.82 K

0.24298 0.00178 0.0002

156.87 ±1.95

0.43562 0.00251 0.0003

146.52 ±3.90

0.56783 0.00343 0.0003 0.95015 0.00617 0.0004 1.02455 0.00635 0.0007 1.52247 0.01011 0.0006 1.55416 0.00984 0.0007 2.03179 0.01359 0.0007 2.04666 0.01296 0.0009 2.55397 0.01773 0.0010 2.56302 0.01646 0.0010 3.09776 0.02132 0.0012 3.08830 0.01974 0.0011

T=347.88 K T=372.86 K

0.66494 0.00558 0.0004

131.11 ±2.84

0.47057 0.00384 0.0003

115.43 ±2.18

1.03324 0.00813 0.0005 0.80554 0.00669 0.0004 1.43453 0.01118 0.0006 1.18951 0.01031 0.0005 1.93326 0.01481 0.0010 1.67869 0.01430 0.0007 2.47450 0.01864 0.0011 2.20515 0.01896 0.0010 3.02742 0.02291 0.0012 2.74701 0.02417 0.0012

Solubility study

89

Figure 4-8. Isothermal phase equilibrium of CO in methyl formate

H2-METHANOL SYSTEM AND H2-METHYL FORMATE SYSTEM

The solubility results of H2 in liquid methanol at different temperatures are given in Table

4-7 and Figure 4-9, and the solubility results of H2 in liquid methyl formate are indicated

in Table 4-8 and Figure 4-10. The solubilities of H2 in liquid methanol and methyl formate

present similar trends with temperature as the solubility of CO in liquid methanol and

methyl formate. The solubilities of H2 in both methanol and methyl formate increase with

increasing the temperature. The solubility of H2 in methanol is lower than that in methyl

formate at the same temperature, resulting in a Henry’s constant obtained in methanol

being greater than in methyl formate. In addition, compared to CO in methanol and

methyl formate at the same temperature, H2 is more difficult to dissolve into both

solvents.

Solubility study

90

Table 4-7. Partial pressure (PH2), liquid phase mole fraction (xi), Henry’s law constant (H) and uncertainties (δ) of H2 in methanol from 298.15 K to 373.15 K

𝐏𝐇𝟐

(MPa) 𝐱𝐢 𝛅

H (MPa)

𝐏𝐇𝟐

(MPa) 𝐱𝐢 𝛅

H (MPa)

T=296.87 K T=320.34 K

0.23685 0.000403 0.00007

588.17 ±3.17

0.26212 0.000497 0.00004

528.65 ±2.5

0.61662 0.00106 0.0001 0.62093 0.00115 0.0001 1.10269 0.00190 0.0002 1.10211 0.00208 0.0001 1.60705 0.00273 0.0002 1.61371 0.00306 0.0002 2.12691 0.00361 0.0004 2.11649 0.00400 0.0003 2.64909 0.00453 0.0004 2.64699 0.00498 0.0005 3.17393 0.00537 0.0005

T=348.26 K T=374.85 K

0.27542 0.000594 0.00008

466.53 ±2.17

0.84095 0.00211 0.0003

397.53 ±1.84

0.57770 0.00124 0.0001 1.02406 0.00257 0.0002 1.04617 0.00223 0.0002 1.47452 0.00368 0.0002 1.56517 0.00334 0.0003 1.98774 0.00500 0.0004 2.09521 0.00452 0.0003 2.52097 0.00638 0.0005 2.60817 0.00560 0.0004 3.04789 0.00766 0.0005

Figure 4-9. Isothermal phase equilibrium of H2 in methanol

Solubility study

91

Table 4-8. Partial pressure (PH2), liquid phase mole fraction (xi), Henry’s law constant (H) and uncertainties (δ) of H2 in methyl formate from 298.15K to 373.15 K

𝐏𝐇𝟐

(MPa) 𝐱𝐢 𝛅

H (MPa)

𝐏𝐇𝟐

(MPa) 𝐱𝐢 𝛅

H (MPa)

T=296.63 K T=325.45 K

0.23819 0.00052 0.00002

480.76 ±10.21

0.44456 0.00117 0.00008

376.54 ±4.40

0.65607 0.00141 0.0001 1.20781 0.00318 0.0001 1.08370 0.00234 0.0002 1.43591 0.00377 0.0002 1.57290 0.00342 0.0002 1.87448 0.00487 0.0002 2.08338 0.00438 0.0003 2.37579 0.00628 0.0003 2.59759 0.00530 0.0004 2.88468 0.00771 0.0004 3.11675 0.00642 0.0004

T=348.38 K T=371.61 K

0.46206 0.00169 0.00006

283.39 ±2.31

0.46720 0.00237 0.0001

187.46 ±6.72

1.12690 0.00395 0.0001 0.65865 0.00321 0.0002 1.43690 0.00509 0.0003 1.09824 0.00593 0.0004 1.84089 0.00656 0.0004 1.53884 0.00842 0.0005 2.30308 0.00814 0.0006 2.08377 0.01104 0.0008 2.82852 0.00992 0.0005

Figure 4-10. Isothermal phase equilibrium of H2 in methyl formate

Solubility study

92

DISCUSSION ON THE SOLUBILITY TREND WITH TEMPERATURE

As discussed in the previous sections, the solubility of both H2 and CO in either methanol

or methyl formate is increased with the increase of temperature. One possible reason is

that the CO-CO and H2-H2 interactions are stronger than CO-solvent and H2-solvent

interactions, and they prefer to associate with themselves rather than the solvents. As the

temperature increases, the solvent undergoes thermal expansion, implying that the

solvent molecules are further apart, and their intermolecular forces are weaker. The gas

molecules may accommodate the solvent molecules more easily. This may increase the

apparent solubilities of CO and H2 gas solutes. However, in conventional vapour-liquid

systems, the gas solute interacts favourably with the liquid and the process is exothermic,

leading to a decrease in gas solubility when the temperature increases.

As discussed earlier, the overall solution process for systems in which solute-solute

interactions are stronger than solute-solvent interactions are endothermic. The binding

energy between methanol and methanol was determined to be 28 kJ/mol [143], however,

the binding energy between CO and methanol and H2 and methanol are not available in

the literature, therefore, water was selected as an example since it did not expect to be

large different from the methanol situation. The binding energy between the two water

molecules is evaluated to be 20 – 25 kJ/mol [144]. Nevertheless, the binding energy of

H2-H2O, and CO-H2O are only 0.71 kJ/mol and 4.05 kJ/mol, respectively [145], [146].

Higher energy input is required to break the water-water bonds than is released when

H2-H2O or CO-H2O bonds are formed, making the process endothermic overall. High

temperatures provide high energy to break solute bonds in the gas phases, facilitating the

dissolving process of the gas solutes with broken bonds into the solvent.

Solubility study

93

THERMODYNAMIC PROPERTIES

The dissolution enthalpy, dissolution entropy and dissolution Gibbs free energy are

shown in Table 4-9. As can be seen, the dissolution enthalpy is positive for all four

systems, indicating that all these processes for gas dissolving into the liquid solvent are

endothermic so that the solubility increases with increasing the temperature. In addition,

with the increase of the temperature, ∆𝑑𝑖𝑠𝐺 increases, implying that CO and H2 dissolved

into methanol and methyl formate are non-spontaneous processes, which requires

external energy to allow the solvent to dissolve into the gas.

Table 4-9. The thermodynamic properties of the systems

System Temperature (K) ∆disH (kJ/mol)

∆disS (J/mol·K)

∆disG (kJ/mol)

CO-methanol 298 4.67 -31.2 14.0 328 14.8 348 15.5 373 16.3

CO-methyl formate

300 3.89 -29.1 12.6 323 13.3 348 14.1 373 14.8

H2-methanol 297 4.57 -37.8 15.8 320 16.7 348 17.7 375 18.7

H2-methyl formate

297 11.2 -14.1 15.4 325 15.8 348 16.1 372 16.5

4.4.3 MODELLING RESULTS

Experimental results on the vapour-liquid equilibrium of CO and H2 in liquid methanol

and methyl formate have been used for the regression of the binary interaction

parameter 𝑘𝑖𝑗 using the phi-phi method in the modified Peng-Robinson Equation of State.

The results for four different systems are given in Table 4-10, and the modelling

Solubility study

94

validation results for CO-methanol, CO-methyl formate, H2-methanol and H2-methyl

formate systems are shown in Figure 4-11, Figure 4-12, Figure 4-13 and Figure 4-14;

respectively. The binary parameter 𝑘𝑖𝑗 is temperature dependent, which can be

described using Equation 4-44, where 𝐴 , 𝐵 and 𝐶 are empirical coefficients [124]. In

some particular vapour-liquid systems, 𝑘𝑖𝑗 may be constant when changing the

temperature [104], and this phenomenon has also been found in CO-methanol and CO-

methyl formate in our study, where the values for these two systems are constant at -0.21

and -0.051. However, for H2-methanol and H2-methyl formate systems, the binary

interaction coefficient is temperature dependent, and the empirical coefficients are

regressed and given in Table 4-11. According to Kordas et al., the binary interaction

parameter 𝑘𝑖𝑗 has different boundary conditions depending on the characteristics of the

solvent [136]. For Non-polar solvent, 𝑘𝑖𝑗 should be located between 0 and 0.2, whereas

for polar solvents, 𝑘𝑖𝑗 should be negative or greater than 0.2. In our study, as methanol

and methyl formate are polar solvents, the regressed binary parameters are either

negative or greater than 0.2 (shown in Table 4-10), which satisfy the boundary conditions

of 𝑘𝑖𝑗 .

The maximum absolute average relative deviation (AARD) are used to check the

reliability of the model for each system. The equation is defined in Equation 4-45, where,

𝑥𝑒𝑒𝑥𝑝 is the experimental data, 𝑥𝑒

𝑐𝑎𝑙𝑐 is the simulation result, 𝑁 is the total number of

experimental points. The results are also listed in Table 4-10, showing that the model can

well represent the experimental data (within 5%), and validation results are shown from

Figure 4-11 to Figure 4-14. Therefore, modified Peng-Robinson Equation of State is an

appropriate model to predict the solubilities of CO and H2 in methanol and methyl

formate.

Solubility study

95

Equation 4-44. The binary interaction parameter

𝑘𝑖𝑗(𝑇) = 𝐴 +𝐵

𝑇+ 𝐶𝑇

Equation 4-45. Definition of AARD

AARD(%) =1

N∑

|xeexp

− xecalc|

xeexp

N

1

× 100

Table 4-10. The regressed binary parameter using PR EoS for different systems

System Temperature (K) 𝑘𝑖𝑗 AARD%

CO-methanol 298.10 -0.21 3.20 322.71 -0.20 0.70 347.76 -0.21 4.30 372.90 -0.21 0.38

CO-methyl formate 299.72 -0.052 0.71 322.82 -0.052 4.70 347.88 -0.051 2.43 372.86 -0.051 2.73

H2-methanol 296.87 -0.33 0.11 320.34 -0.24 0.30 348.26 -0.12 0.95 374.85 -0.015 0.19

H2-methyl formate 296.63 0.36 3.70 325.45 0.40 1.70 348.38 0.39 0.40 371.61 0.28 3.50

Solubility study

96

Figure 4-11. Modelling validation results of CO solubility in methanol

Figure 4-12. Modelling validation results of CO solubility in methyl foramte

Solubility study

97

Figure 4-13. Modelling validation results of H2 solubility in methanol

Figure 4-14. Modelling validation results of H2 solubility in methyl formate

Solubility study

98

Table 4-11. Empirical coefficients of binary interaction parameters 𝑘𝑖𝑗

System 𝑨 𝑩 𝑪 H2-methanol -1.74 33.99 0.000437 H2-methyl formate 12.63 -1973.1 -0.019

4.5 CONCLUSIONS

In this chapter, the vapour-liquid equilibrium of CO and H2 in liquid methanol and methyl

formate was measured in a custom designed apparatus from 296 K to 375 K. The

apparatus was pre-validated using published CO2 and CO solubility data in methanol. It is

found that the solubilities of CO in methanol and methyl formate are greater than those

of H2 in methanol and methyl formate, and both CO and H2 solubilities in these two liquid

solvents increase with the increase of the temperature. This leads to an endothermic

process. The experimental results were used for the regression of modified Peng-

Robinson Equation of State (PR EoS), and the binary interaction parameter 𝑘𝑖𝑗 were

determined via a phi-phi method. The binary parameters for CO-methanol and CO-methyl

formate system are temperature independent, whereas the binary parameters for H2-

methanol and H2-methyl formate system are temperature dependent. All the binary

parameters satisfy the boundary constraint conditions due to the polar characteristics of

methanol and methyl formate. Through employing the PR EoS model, the solubility can

be predicted well for the investigated system.

Hydrogenation reaction kinetics mechanism

99

CHAPTER 5 HYDROGENATION REACTION KINETICS

MECHANISM

5.1 INTRODUCTION

In the methanol synthesis via the methyl formate (MS via MF) two-step approach, the

hydrogenation reaction (the second step) is the rate limiting step. However, to our best

knowledge, few studies have been reported on the hydrogenation mechanisms and the

characteristics of the catalysts. Hence, it is important to investigate the hydrogenation

reaction to understand comprehensively the reaction mechanism and kinetics. In

addition, since the first step is commercialised in ester production, its kinetics are well

known. Study of the second step becomes crucial and a clear understanding may aid in

the scale up of the production of methanol via the ‘MS via MF’ approach at moderate-

temperature and moderate-pressure conditions.

Producing alcohols from esters is one of the fundamental reactions in organic chemistry,

and has been employed in a great number of chemical manufacturing industries [147]. In

industry, Adkin-type catalysts (CuO/CuCr2O4) are used to produce alcohols due to their

stable structure although they are operated at harsh conditions such as temperatures

ranging from 200 to 300 °C and H2 pressures ranging from 14 to 30 MPa [147], [148].

However, side reactions and degradation of the reactants and products are the main

problems in the current manufacturing processes [149].

In our present work, methanol is produced from hydrogenation of methyl formate, which

is given by Reaction 5-1. In this reaction, catalytic hydrogenation of carboxylic esters is

challenging due to its low electrophilicity of carbonyl carbon and the difficulties

associated with polarizing the carbonyl group of the substrate [150]. A number of


100

researchers have studied a range of hydrogenation catalysts intensively, including

heterogenous catalysts and homogeneous catalysts, in order to screen cost effective and

environmentally friendly catalysts [148], [149], [151].

Reaction 5-1. Hydrogenation of methyl formate

CH3COOH + 2H2 ⇌ 2CH3OH

Regarding the homogenous catalysts which have recently been published in the literature,

Ru- and Os-based homogenous catalysts were developed for the hydrogenation of esters

but were only limited to highly activated esters [152]. Significant breakthroughs were

made in the catalytic hydrogenation of esters by Milstein et al. who developed Ru-based

catalysts for the hydrogenation of non-active esters in 2006 [11]. Later on, Pidko group

proposed a new pathway for the hydrogenation of methyl formate to produce methanol

using bipyridine-based Ru-pincer complex at 110 °C and 5 MPa [151]. However, such

catalysts contained expensive metals and toxic chemicals, which makes it commercially

unattractive.

In comparison with the homogeneous catalysts, heterogeneous catalysts, such as copper

chromite and copper-based catalysts, were widely investigated in both gas and liquid

phases [64], [155], [156]. The active sites of such copper-based catalysts for ester

hydrogenation reactions are considered as Cu0, Cu+, or a combination of Cu0 and Cu+. The

heterogeneous catalysts are more favourable as they have improved compatibility with

easy recovery. The solid-state heterogeneous catalysts are also easy to separate and

recover after finishing experiments. Copper chromite is one of the most commonly used

catalysts to evaluate the hydrogenation of esters to produce alcohols. The copper

chromite is manufactured via the discompose of either copper barium ammonium

chromite (Ba2Cu2(NH4)2(CrO4)5) or copper ammonium chromite (Cu(NH4)2(CrO4)2) at a


101

temperature range from 350 °C to 450 °C (shown in Reaction 5-2 and Reaction 5-3,

respectively). In addition, a separation process of the reaction products is required to

purify the final products.

Reaction 5-2. Decomposition of copper barium ammonium chromite

Ba2Cu2(NH4)2(CrO4)5 ⇌ CrCuO3 + CuO + 2Ba + 4H2O + 4 Cr + N2 + 6O2

Reaction 5-3. Decomposition of copper ammonium chromite

Cu(NH4)2(CrO4)2 ⇌ CrCuO3 + CrO + 4H2O + N2

Since the copper chromite has good compatibility, stability and is easy to recover, copper

chromite is employed in the current research to study the reaction kinetics and the

possible mechanism. In this chapter, three objectives were discussed and explained.

Firstly, to study the carbonylation reaction to verify the reaction and compare with

literature data; secondly, to study the structure, surface properties and thermal stability

of the catalyst by using XRD, SEM, TGA, N2 physisorption, TPR and N2O chemisorption

techniques; Finally, to investigate the kinetics and the possible reaction mechanism of

copper chromite catalysed hydrogenation of methyl formate. A number of experiments

were performed in a batch reactor at different temperatures and pressures ranging from

346 to 384 K and 1.8 to 2.2 MPa, respectively, with a range of copper chromite

concentrations (8 g/L to 20 g/L).


102


5.2.1 APPARATUS

Figure 5-1. Schematic diagram of reaction apparatus: 1. Gas cylinders (He, CO2, H2 and CO); 2. Mass flow controller, 3. Storage tank; 4. Reactor; 5. Heating tape/Cooling bath; 6. Magnetic stirrer; 7. Vacuum pump; 8. Vent system; 9. Gas sampling tank; 10. GC; BV-1 to BV-6: Ball valves; NV-1 to NV-3: Needle valves

The bench scale reactor system was designed and set up in the lab to perform the

reactions. The gases (including He, CO and H2) were charged from gas cylinder 1 into the

storage tank 3 by manipulating mass flow controllers 2 for each gas specifically. The

storage tank 3 serves the purpose of the determination of the amount of gases fed to the

reactor 4 by evaluating the pressure difference of the storage tank between before and

after the gas loading. The reactions were conducted in the reactor 4 and equipped with a

magnetic stirrer 6 to thoroughly mix the reactants and the catalysts. The vacuum pump 7

was used to remove the air in the storage tank 3 and reactor 4 before initiating the

reactions. The gas products were collected by a gas sampling tank 9 and analysed by a

gas chromatography 10.

Pressures of storage tank 3 and reactor 4 were monitored with a Swagelok S-model

pressure transducer (accuracy ≤ 0.25% span limit point calibration) and a MKS Type

TT A

1

6

BV-1 BV-2 BV-3 BV-4

NV-1 NV-2

2

3

PT A PT B TT B

4

5

8 8

Computer

7

BV-6

Bv-7

BV-5

9

10


103

627B pressure transducer (accuracy ± 1 mmHg), respectively. Both pressure transducers

were connected to a NI™ USB-6002 data acquisition (DAQ) device and pressure values

were recorded by LabVIEW System Design Software to record the pressures. Meanwhile,

the pressure of the reactor was real-time displayed by a power supply digital readout unit

(MKS Instruments Inc, 660B model). Two pressure transducers were calibrated with

accurate manometers to provide accurate and reproducible results.

The temperature inside the reactor 4 was controlled with an external heating tape

coupled with a temperature controller 5. Temperatures of both storage tank 3 and

reactor 4 were measured with K-type thermocouples (accuracy ± 0.1 K) connected to

National Instruments™ (NI™) thermocouple measurement devices and recorded with

LabVIEW System Design Software.

5.2.2 PROCEDURE

The procedure of methanol synthesis at low temperature and low pressure comprises the

following steps:

1) Leak Test: Prior to every experiment, the leak test was undertaken by injecting helium

into the reactor (4) to 30 bar and monitoring the pressure stability. The leak test takes

two hours.

2) Loading catalysts and liquid reactant and system degassing: About 12 mL of liquid

reactant (methanol or methyl formate) and catalysts were loaded to the reactor (4). The

reactor was subsequently connected to the process apparatus line via a VCR-gasket

(Swagelok®), and then immersed in dry ice bath to cool down. The cooling step by dry

ice is to reduce the vapour pressure of the volatile solvents used in the experiments of

this work to an acceptable value (less than 500 Pa). Once the temperature inside of the


104

reactor reduced below 213 K and the vapor pressure of the solvents reaches to below 200

Pa, a vacuum pump was turned on to degas the reactor. The solvent loss during the degas

step was less than 0.5 wt % as was calculated by the mass balance, and it is assumed that

there was no other gas in the reactor after the degassing procedure.

3) Gas injection to storage tank: The valves BV-1 and BV-2 were opened to introduce the

H2 or CO from gas cylinder 1 via mass flow controllers to the storage tank 3 to a

predetermined value.

4) Catalytic reaction step: The gas feed was charged into the reactor 4 from the storage

tank 3 by opening ball valve (BV-3) for 3 only seconds, this allows the gas to flow to the

reactor and eliminates the back flow to tank 3. Following this, the heating tape was

attached and set to a reaction temperature. Once the temperature reached to the desired

value and the pressure in tank 4 stabilized, the stirrer was turned on to thoroughly mix

the gas and the liquid and hence initialize the reaction. Simultaneously, the pressure and

temperature values of both tanks were recoded every 15 seconds via the LabVIEW

software. The reaction was considered at an equilibrium state when the pressure inside

the reactor (tank 4) stabilized for two hours. In addition, it was found that an induction

period of 12 hours and one hour are present at low temperature at 346 K and 370 K,

respectively. Such an induction period is absent at high temperatures of 384 K.

5) Liquid and gas sampling: Due to the harsh operating conditions of high pressure and

high temperature, regular sample collection during the experiment for analysis purposes

was very difficult. Therefore, the experiment was regularly terminated to allow us to

collect sample and investigate the reaction progress. The regular sampling interval was

selected based on the overall time for the reaction to reach equilibrium. To terminate the

experiment and take out a liquid sample, the heating tape was quickly removed and the


105

reactor (tank 4) was quickly quenched by dry ice to cool down to 213 K. This would

terminate the reaction and prohibit the liquid sample vaporization during the pressure

relief process. The gas products were collected by using a gas sampling tank 9 which is

connected to the valve BV-5. By turning on the valves (BV-4 and BV-5), gas moves from

the system to the gas sampling tank. The liquid products were taken out from the reactor

4.

6) The sample analysis: Both liquid and gas samples were analysed via a gas

chromatography (GC).

7) Catalyst recovery: The solid catalysts can be recovered by filtering the liquid products,

and then washed with the liquid reactant. After that, the solid catalysts were vacuum

dried at room temperature for 12 hours.

In this chapter, experiments were initially carried out at different rotation speeds to

determine the optimized agitation speed with minimum mass transfer resistance. The

effect of catalyst loadings was also investigated. Subsequently, experiments were

conducted at three different temperatures using the optimized agitation speed and

catalyst loading to determine the activation energy (𝐸𝑎) of the copper chromite catalysed

hydrogenation reaction. A possible reaction mechanism was proposed and validated by

running a series of experiments at different hydrogen pressures. The experimental

conditions of the hydrogenation reaction are listed in Table 5-1.

Table 5-1. Experimental conditions for hydrogenation reactions

Parameters Experimental conditions Catalysts Copper chromite Rotation speeds (rpm) 400, 600, 800 and 1000 Catalyst loadings (g/L) 8, 12, 16 and 20 Temperatures (K) 346, 370 and 384 Pressure (MPa) 1.8, 2.0 and 2.2


106

5.3 CARBONYLATION REACTION STUDY

In order to confirm the literature results of the carbonylation reaction, three validation

experiments were conducted at 50 °C, 75 °C and 100 °C and their total pressure profiles

are summarized in Figure 5-2. Potassium methoxide was employed as the catalyst at the

concentration of 0.4 mol/L. Pure carbon monoxide and methanol were fed into the

reactor. The time to reach equilibrium and the equilibrium conversions are tabulated in

Table 5-2. The pressure profiles show that the reaction rates increased with the

increasing temperatures, while the equilibrium conversions decreased with the

increasing temperatures. In industry, the carbonylation of methanol is carried out at

approximately 70 to 75 °C which compromises the reaction rate with the equilibrium

conversions [157]. Compared with the hydrogenation, the operating temperature of the

carbonylation reaction was much lower than that of the hydrogenation reaction. Hence,

it is confirmed that the hydrogenation reaction is the rate-determining step in the system.


107

Figure 5-2. The pressure profiles of carbonylation reaction. Operating conditions: Ptotal=2.3 MPa, agitation speed = 800 rpm, catalyst loadings = 0.4 mol/L

Table 5-2. The carbonylation reaction performance

Temperatures(°C) Equilibrium conversions (%) Time to reach equilibrium 50 90.5 80 minutes 75 59.8 11.67 hours 100 27.3 56.38 hours

5.4 HYDROGENATION REACTION CATALYSTS PREPARATION AND

CHARACTERISATION

5.4.1 CATALYSTS PREPARATION

Copper chromite is reduced in a hydrogen atmosphere before conducting the

hydrogenation reaction, since it is necessary to transform the copper oxide phase in the

catalyst into metallic copper which represents the active site in the catalysts. Therefore,

it is important to understand its reducibility to determine the appropriate reduction

temperature. The temperature-programmed reduction (TPR) measurement was

conducted.

The reduction profile (H2 consumption) of copper chromite is shown in Figure 5-3. It

exhibits a broad reduction peak accompanied by a shoulder of 150 – 230 °C. Accordingly,

to achieve a complete reduction, it is important to select a reduction temperature which

is higher than 230 °C thus a temperature of 260 °C was selected in this work.


108

Figure 5-3. The TPR profile of the copper chromite catalysts

In addition, the TPR profile is deconvoluted into two Gaussian peaks. Both peaks and their

corresponding contributions are listed in Table 5-3. In literature, the peak (peak α) at low

temperature is attributed to the reduction of dispersed CuO, whereas the peak at high

temperature (peak β) indicates the reduction of CuO in bulk phase which has a larger size

compared to the dispersed CuO [158]. As can be seen in Table 5-3, 39.68 % of CuO in the

copper chromite is in a dispersed form.

Table 5-3. The reducibility of the copper chromite

Sample TPR peak position (°C) and concentration (%)a H2 consumption (mmol/g) Peak α Peak β

Copper Chromite 176 (39.68) 197(60.32) 6.282 a Values in parentheses are the contributions (%) of each peak

5.4.2 THE STRUCTURE AND PHASE COMPOSITIONS

To identify the structure and the phase composition of copper chromite, the powder X-

ray diffraction technique is employed. The XRD patterns of raw and reduced copper

chromite are shown in Figure 5-4. Pattern (a) shows four diffraction peaks, each belongs

to a different species. The sharp diffraction peak at 2θ = 35.4° and 38.8° are ascribed to


109

the crystal copper oxide (CuO), (JCPDS card #45-0937), which is the main form of copper.

Two typical diffraction peaks at 2θ = 42.3° and 64.8° are attributed to (2 2 0) and (4 1 1)

crystal planes of a crystalline copper chromium oxide (CuCrO4)[159]. The other small

peaks appeared on the diffraction pattern suggest the presence of traces of crystalline

mcconnellite (CuCrO2) and chromium (III) oxide (Cr2O3) [159], [160].

The XRD pattern of the reduced copper chromite using pure H2 treatment at 533 K for 6

hours is shown in Figure 5-4 (b). The pattern exhibits three typical peaks at 2θ = 43.3°,

50.4° and 74.1°, which are assigned to (1 1 1), (2 0 0) and (2 2 0) crystal planes of a

crystalline copper phase (JCPDC card #04-0836). After reduction, the diffraction peaks

attributed to CuO disappeared, but the peaks of well crystallised copper phase are

present, suggesting that the copper species are thoroughly reduced during the hydrogen

reduction. In addition, the peaks assigned to Cr2O3 disappeared; instead, new peaks

ascribed to chromium (IV) oxide (CrO2) appeared [159]. The comparison between the

characteristics of copper chromite before and after the hydrogen reduction indicates that

the intensity of crystalline CuCrO4 phase is decreased but can still be detected.


110

Figure 5-4. XRD patterns of copper chromite catalyst

5.4.3 THERMAL BEHAVIOUR AND STABILITY

The thermal gravity analysis (TGA) was conducted to study the thermal behaviour of the

catalyst before and after reduction. The TGA diagram as well as the differential

thermogravimetric (DTG) curve shown in Figure 5-5 indicate that the total weight loss

(%) of the sample is 12.58 %. According to the TGA results, the catalysts showed two

major weight loss stages during the thermal decomposition process, which are at 290 °C

and 730 °C. The first weight loss step occurred at 290 °C is due to loss of moisture and

physically adsorbed water. The second weight loss stage at 730 °C is due to the

degradation of the catalyst and the structure change [161].


111

Figure 5-5. The profile of thermal gravity analysis and the corresponding DTG

5.4.4 THE MORPHOLOGY AND SIZE

The catalyst morphology and size of catalysts can be determined using field emission

scanning electron microscopy (SEM). The SEM images of the raw and the reduced copper

chromite are shown in Figure 5-6. The low magnification image of the catalyst, (Figure 5-

3(a)), shows that it consists of irregular shaped chunks that have no distinct morphology.

At high magnification (Figure 5-3 (b)), on the other hand, two different structures are

identified; nanorods (in the green squares) with the length of 200 to 300 nm and a

diameter of 30 nm which might be chromium (III) oxide [127] and flaky aggregates (in

the red squares) consisting tiny particles (about 10 – 20 nm ID) which might be the other

component in the catalyst, copper oxides. However, the SEM images of the reduced

catalyst show two different structures: namely spicula shape structures (in the orange

squares) which might be chromium (II) oxides and hexagonal-rhombohedral structures

(in the blue circles) with size of 600 - 700 nm (Figure 5-6 c & d) [127]. The XRD pattern


112

of the reduced catalyst shows three defined peaks which are attributed to copper crystals

(Figure 5-4). From this, it can be inferred that the hexagonal-rhombohedral structures

are the metallic copper crystals. Moreover, after reduction the nanorod crystals were

disappeared and replaced with spicula shape structures with 300 nm length and 20 nm

diameter indicating that Cr2O3 is transformed into CrO2. In conclusions, the XRD and SEM

result indicate that the active species of metallic copper is successfully reduced and ready

to use for the hydrogenation reaction.

Figure 5-6. SEM images of the copper chromite sample. (a) and (b) are copper chromite; (c) and (d) are reduced copper chromite

5.4.5 THE SURFACE AREA AND SPECIFIC COPPER SURFACE AREA

The ASAP2010 was used to determine the specific surface area and pore volume of

copper chromite, and the results of before and after hydrogen reduction are listed in

1 µm

(a)

0.1 µm

(b)

1 µm

(c)

0.1 µm

(d)


113

Table 5-4. It was found that the surface area and the pore volume are higher after

reduction which might be due to the increase of exposed metal copper content. The BET

(Brunauer-Emmett-Teller) specific surface area of the copper chromite sample used for

the hydrogenation reaction experiments is 136.5 m2/g, This contradicts with Monti’s

result, who claimed that BET surface area of copper chromite after hydrogen reduction

was only 24.8 m2/g, whereas Sorum and Onsager’s work, where BET result was 116 m2/g

[65]. The dispersion of copper (DCu) and the exposed copper surface area (SCu) were

measured via the adsorption of reactive N2O. The values were calculated based on the

information provided in section 3.2.2.2.4 and data are given in Table 5-4.

Table 5-4. Surface properties of the copper chromite

Sample BET specific surface area

(m2/g)

Pore volume (cm3/g)

Cu surface area b

(m2/g)

Cu dispersionb

(%) Before reduction

After reduction a

Before reduction

After reduction a

Copper chromite 41.53 136.5 0.1590 0.2120 51.62 39.90 a After reduction at 260 °C. b Calculated from N2O dissociative adsorption.

5.4.6 SUMMARY

The properties of the copper chromite were investigated by using the TPR, TGA, SEM and

BET surface area. It is found that the reduction temperature is required to be above

230 °C and the BET surface area after reduction is 136.5 m2/g. Hexagonal crystal metallic

copper is produced after the reduction as shown in the SEM photos.

5.5 STUDY OF HYDROGENATION REACTION KINETICS AND EXPLORE THE

REACTION MECHANISM

5.5.1 EFFECTS OF AGITATION SPEED ON REACTION RATE

The hydrogenation reaction of methyl formate to methanol is a gas-liquid-solid reaction

or so called three phase catalytic reaction. The reaction rate may be affected by both


114

external and internal mass transfer resistance which need to be eliminated by a careful

selection of stirring speeds and catalysts particle size.

The production of methanol over the reaction time at various stirring speeds was studied

to verify the effect of stirring speeds on the external mass. Figure 5-7 provides the results

of experiments where identical amounts of pure methyl formate and copper chromite

catalysts were used under different agitation speeds at 384 K and an initial total pressure

of 3 MPa.

Figure 5-7. Effect of rotation speeds on the conversion of methanol. Operating conditions: Ptotal = 3.2 MPa, T = 384 K, Catalyst loading = 16 g/L

As can be seen in Figure 5-7, the initial production rates of methanol are increased with

the increase of agitation speeds from 400 rpm to 800 rpm. The data of the methanol

produced over time was supplied in Table 5-5. It is also found that the production rate of

methanol almost overlapped at 800 rpm and 1000 rpm. In addition, the production rate

at lower rotation speeds approached to those at higher rotation speeds by the end of the


115

experiment, implying that the influence of the external diffusion resistance can be

neglected when the agitation speed is selected over 800 rpm and the reaction is not

controlled by the mass transfer. Therefore, all hydrogenation reaction experiments were

carried out at an agitation speed of 800 rpm.

Table 5-5. Amount of methanol produced under different agitation speeds

400 rpm 600 rpm Time (min) Moles of methanol Time (min) Moles of methanol 0 0 0 0 40 ± 2.5 0.00876 30 ± 2.5 0.0081 80 ± 2.5 0.0138 60 ± 2.5 0.0142 120 ± 2.5 0.0176 90 ± 2.5 0.0182 160 ± 2.5 0.0189 120 ± 2.5 0.0204 200 ± 2.5 0.0198 150 ± 2.5 0.0209 240 ± 2.5 0.0205 180 ± 2.5 0.0209 280 ± 2.5 0.0210

800 rpm 1000 rpm Time (min) Moles of methanol Time (min) Moles of methanol 0 0 0 0 30 ± 2.5 0.00993 30 ± 2.5 0.00962 60 ± 2.5 0.0150 60 ± 2.5 0.0153 90 ± 2.5 0.0189 90 ± 2.5 0.0191 120 ± 2.5 0.0203 120 ± 2.5 0.0205 150 ± 2.5 0.0207 150 ± 2.5 0.0208

Moreover, the internal diffusion can be neglected when the average particle size of the

catalysts is less than 50 µm [162]. Based on the SEM results shown in Figure 5-6, the

active species copper of the copper chromite catalysts has a size of 200 - 300 nm, which

indicates that the internal mass-transfer resistance can be neglected in our experiments.

5.5.2 EFFECTS OF CATALYSTS LOADINGS ON REACTION RATE

The catalysts provide an alternative route for the reaction with a lower activation energy,

which speeds up the reaction. Hence, it is necessary to evaluate the effects of catalyst

loadings on the reaction rate.


116

Figure 5-8. Effect of the catalyst loadings on the conversion rate of methanol. Operating conditions: Ptotal= 3.2 MPa, T = 384K. Rotation speed = 800 rpm.

Figure 5-8 represents the effect of catalyst loadings range between 8 g/L and 20 g/L on

hydrogen conversion. The data are summarised in Table 5-6. The unit of the catalyst

loadings represents the gram of catalysts per volume of methyl formate used in the

experiment. The result indicates that 8 g/L of the reduced copper chromite gives the

lowest methanol production rate. In addition, the hydrogen conversion rate can be

significantly improved by increasing the catalysts loading from 8 g/L to 12 g/L, which

implies that the reaction rate is proportional to the catalysts loadings at the investigated

range. Interestingly, it was found that the reaction rate can only be improved slightly

when adding the catalysts from 12 g/L to 20 g/L, which may suggest that the active sites

of the catalysts should be sufficient to initiate the hydrogen reaction. Therefore, 16 g/L

of catalysts concentration was used for our all hydrogenation reactions.


117

Table 5-6. Amount of hydrogen at various catalyst loadings

8 g/L 12 g/L Time (min) Moles of hydrogen Time (min) Moles of hydrogen 0 0.0490 0 0.0492 40 ± 2.5 0.0417 30 ± 2.5 0.0389 80 ± 2.5 0.0358 60 ± 2.5 0.0328 120 ± 2.5 0.0317 90 ± 2.5 0.0287 160 ± 2.5 0.0290 120 ± 2.5 0.0272 200 ± 2.5 0.0272 150 ± 2.5 0.0264 240 ± 2.5 0.0264

16 g/L 20 g/L Time (min) Moles of hydrogen Time (min) Moles of hydrogen 0 0.0490 0 0.0486 30 ± 2.5 0.0393 30 ± 2.5 0.0386 60 ± 2.5 0.0327 60 ± 2.5 0.0310 90 ± 2.5 0.0283 90 ± 2.5 0.0276 120 ± 2.5 0.0267 120 ± 2.5 0.0266 150 ± 2.5 0.0264 150 ± 2.5 0.0264

5.5.3 EFFECTS OF TEMPERATURE ON REACTION CONVERSION AND SELECTIVITY

To check the temperature effect on the hydrogenation reaction, the experiments were

conducted at 346 K, 370 K and 384 K. The change of the amount of reactants and products

over reaction time at these different temperatures were shown in Figure 5-9 to Figure

5-11, and their corresponding hydrogen conversion ratios were also given in these

figures. The experimental data are summarised in Table 5-7. The results indicate that

hydrogen conversion increases from 46.12 % to 83.68 % when the temperature

decreases from 384 K to 346 K. This is because the hydrogenation reaction is an

exothermic reaction, which thus prefers at low operating temperatures. However, the

reaction time required to reach equilibrium state significantly decreased from

approximate 140 hours to 150 minutes when the reaction temperature increases from

346 K to 384 K. This is because high temperatures accelerate the hydrogenation reaction

due to the increase the Brownian movements of reactant molecules, which increases the

opportunities to incur the reaction.


118

Figure 5-9. Effect of temperature on the reaction rate. Operating conditions: Ptotal = 3.2 MPa, T = 346 K, Catalyst loading = 16 g/L. Rotation speed = 800 rpm.


119



120



121

Table 5-7. Experimental values of reactants and products at different temperature.

346 K Time (hours)

Moles of H2 (mole)

Moles of CH3OCOH (mole)

Moles of CH3OH (mole)

Conversion of H2

(%) 0 0.0647 0.183 0 0 22 ± 0.1 0.0555 0.170 0.00908 14.3 44 ± 0.1 0.0464 0.157 0.0180 28.3 66 ± 0.1 0.0366 0.143 0.0276 43.43 88 ± 0.1 0.0263 0.128 0.0376 59.35 110 ± 0.1 0.0168 0.115 0.0469 74.01 132 ± 0.1 0.0112 0.107 0.0524 82.75 138.3 ± 0.1 0.0106 0.106 0.0530 83.68

370 K Time (hours)

Moles of H2 (mole)



Conversion of H2

(%) 0 0.0560 0.173 0 0 4 ± 0.2 0.0434 0.153 0.0146 22.63 8 ± 0.2 0.0345 0.139 0.0230 38.41 12 ± 0.2 0.0277 0.129 0.0295 50.56 16 ± 0.2 0.0239 0.123 0.0332 57.43 20 ± 0.2 0.0216 0.119 0.0353 61.50 24 ± 0.2 0.0201 0.117 0.0366 64.18 28 ± 0.2 0.0195 0.116 0.0373 65.22 32 ± 0.2 0.0188 0.115 0.0380 66.52 36 ± 0.2 0.0183 0.114 0.0384 67.34 40 ± 0.2 0.0182 0.114 0.0385 67.45

384 K Time (minutes)

Moles of H2 (mole)



Conversion of H2

(%) 0 0.0490 0.171 0 0 30 ± 2.5 0.0393 0.153 0.00925 19.86 60 ± 2.5 0.0327 0.141 0.0155 33.27 90 ± 2.5 0.0283 0.133 0.0195 42.24 120 ± 2.5 0.0267 0.130 0.021 45.51 150 ± 2.5 0.0264 0.130 0.021 46.12


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The analysis of reaction products show trace amounts of undesired products, CO and

dimethyl ether in the gas phase of some experiments, which can be attributed to a

decarbonylation reaction (shown in Reaction 5-4) and a methanol dehydration reaction

(shown in Reaction 5-5), respectively. As their produced amount are small, it is difficult

to evaluate their peak areas from the GC analysis, and thus a 99 % selectivity of the

hydrogenation reaction may be assumed in the operating temperatures ranges between

346 K and 384 K. On the other hand, it is noted that a large portion of methyl formate was

detected in the gas phase due to its high vapour pressure (c.a. 1 MPa) at 384 K.

Reaction 5-4. Decarbonylation reaction

HCOOCH3 ⇌ CH3OH + CO

Reaction 5-5. Dehydration of methanol

2CH3OH ⇌ CH3OCH3 + H2O

It can be found that the concentration of reactants decreases dramatically at the initial

stage of the reaction and then slows down gradually with increasing the reaction time.

Therefore, most of the hydrogen reaction occurs at the beginning of the experiments due

to high concentration of reactants.

5.5.4 REACTION KINETICS MODEL AND PARAMETER ESTIMATION

In order to obtain the reaction rate, it is necessary to consider a reaction mechanism.

However, no published literature on the possible hydrogenation reaction mechanism is

available, it is important to postulate a reaction mechanism for the current hydrogenation.

Hence, a kinetics model with the dissociative adsorption of hydrogen is derived. In the

theory, the molecular size of hydrogen is much smaller than the organic species, and thus

it can be adsorbed on the surface of the solid catalysts freely. Therefore, the adsorption

of hydrogen and methyl formate on the catalysts can be considered as non-competitive.


123

The elementary reactions of this mechanism can be found in Reaction 5-6 to Reaction

5-10.

Reaction 5-6. Adsorption of methyl formate on the catalyst active sites

HCOOCH3 + 2s ⇌ HCOOCH3(s2)

Reaction 5-7. Adsorption of H2 on the catalyst active sites

2H2 + 4s ⇌ 4H(s)

Reaction 5-8. Formation of intermediates

HCOOCH3(s2) + 2H(s) ⇌ CH2OH(s) + CH3O(s) + 2s

Reaction 5-9. Production of methanol I

CH2OH(s) + H(s) ⇌ CH3OH + 2s

Reaction 5-10. Production of methanol II

CH3O(s) + H(s) ⇌ CH3OH + 2s

The reaction mechanism involves a major heterogeneous reaction where the active

reactant molecules are adsorbed onto the surface of solid catalysts. In this step, strong

chemical bonds are formed between reactant molecules and catalysts during the

adsorption process, which are considered as reversible reactions. The active sites on the

surface of the copper chromite catalyst is denoted as ‘s’ in this chapter. In the current

system, methyl formate and hydrogen are the two reactants, and their molecules are

taken into account. In the proposed mechanism, methyl formate presents as the methyl

formate molecules which can be attached onto the active sites of copper chromite

catalysts (Reaction 5-6), while hydrogen molecule (dihydrogen) is dissociated as two

chemisorbed hydrogen atoms which can be bonded with the catalytic active sites

(Reaction 5-7). In addition, H2 molecules can be adsorbed dissociatively to occupy two

surface sites, hence, the rate of adsorption depends on the number of pairs of available


124

surface sites. The reserve reaction is known as associative desorption. Two intermediates

(Reaction 5-8), CH2OH(s) and CH3OH(s) are subsequently generated followed by

Reaction 5-6 and Reaction 5-7, which becomes the precursor reaction to produce the

methanol molecules (Reaction 5-9 and Reaction 5-10).

In literature, the quasi equilibrium method is widely applied to determine reaction

kinetics models and their corresponding rate expressions. In this model, several reaction

mechanisms are proposed, where there are many elementary steps. For example,

hydrogenation reactions contain different carbon chain repartition and monomer

formation elementary steps. In each proposed reaction mechanism, only one elementary

step is assumed to be a rate limiting step (RDS), and all other steps are considered as

equilibrium, or named quasi equilibrium. The expression is constrained with the site

balance where the sum of surface fraction adsorbed by different species equals to one.

Based on the assumption, a rate equation can be expressed.

In each mechanism, every elementary step has possibilities to be the RDS, thus having a

number of rate expressions for different RDS in different mechanisms. All the expressions

are further required to be regressed using experimental data to evaluate the parameters

in each expression. The one with the best fitting is considered the most likely reaction

mechanism and reaction kinetics model. According to Mirzaei et al., there are many ways,

including graph, residual plot, confidence interval, 𝑅2 , 𝑅𝑎𝑑𝑗2 and variance and 𝑅𝑚𝑠𝑑 , to

evaluate if the rate expression reflects the experiment correctly [163]. The graph is used

to track the reaction rate of reactants or products to check if the prediction of the model

is consistent with the experiment. The residual plot is used to check the deviation

between the model and the experiment. Confidence intervals are used to check the

robustness of evaluated parameters in the model. 𝑅2 and 𝑅𝑎𝑑𝑗2 are employed if the model


125

responds to the experimental data. Variance and 𝑅𝑚𝑠𝑑 are adopted to evaluate the

accuracy of the model.

It was found that using the quasi equilibrium method is a time-consuming process, which

involved great amounts of work in deriving rate expressions and evaluating the best

model. In addition, some reactions may have multiple rate limiting steps, and other

elementary reactions may not be at equilibrium. Therefore, using this method may not be

reliable and accurate to illustrate a reaction mechanism.

Hence, in our study, a new approach will be used and presented. Since the elementary

reaction steps are proposed, the reaction rate of each component including the

intermediates can be expressed based on the concentration and the kinetics constants,

and all those equations are ordinary differential equations (ODEs). The ODEs can be

found from Reaction 5-6 to Reaction 5-10. In MATLAB, such ODEs can be solved using the

built-in MATLAB solver ‘ode15s’. ‘ode15s’ adopts variable-step, variable-order based

numerical differentiation formula (NDF) and backward differentiation formula (BDF) to

solve ODEs in the form of Taylor polynomial. The variable order for Taylor polynomial

can be selected up to 5 in ‘ode15s’ to evaluate the objective function 𝜕𝐹

𝜕𝑡. A typical BDF

using higher orders of Taylor polynomial is given by Reaction 5-9. Considering the

concentration of each element in the proposed reaction mechanism is in a broad range,

which make the objective function stiff, the numerical method should take small variable

steps to obtain satisfactory results. Hence, the ‘ode15s’ provides an efficient and time-

saving method without sacrificing the calculation accuracy.

The kinetic parameters 𝜃 = [𝑘1, 𝑘−1, 𝑘2, 𝑘−2, 𝑘3, 𝑘−3, 𝑘4, 𝑘−4, 𝑘5, 𝑘−5] were evaluated

using the least-squares regression method, and the objective function is given in Equation

5-10, which is the sum of difference between the concentration of hydrogen, methyl


126

formate and methanol in experiments and in kinetics models. A built-in function ‘fmincon’

was employed to determine the parameter values in MATLAB R2017b software. ‘fmincon’

is a common method using interior point algorithm to optimize a group of non-linear

equations and find out the minimum value of the objective function. The initial values

were given based on the amount of reactants fed into the reactor. Due to the reversible

reactions in Reaction 5-6 and Reaction 5-7, a minimum value of 1 × 10−5 is set for the

products, [𝐻𝐶𝑂𝑂𝐶𝐻3(𝑠2)], [𝐻(𝑠)], [𝐶𝐻3𝑂(𝑠)] and [𝐶𝐻2𝑂𝐻(𝑠)] in order to solve the ODEs

in the MATLAB. An initial value of 0.04 was suggested for [𝑠], and can be further tuned to

fit the experimental data.

Equation 5-1. Reaction rate expression of methyl formate

d[HCOOCH3]

dt= −k1[HCOOCH3][s]2 + k−1[HCOOCH3(s2)]

Equation 5-2. Reaction rate expression of H2

d[H2]

dt= k2[H2]2[s]4 − k−2[H(s)]4

Equation 5-3.Reaction rate expression of methanol

d[CH3OH]

dt= k4[CH2OH(s)][H(s)] − k−4[CH3OH][s]2 + k5[CH3O(s)][H(s)] − k−5[CH3OH][s]2

Equation 5-4. Reaction rate expression of HCOOCH3(s2)

d[HCOOCH3(s2)]

dt

= k1[HCOOCH3][s2] − k−1[HCOOCH3(s2)] − k3[HCOOCH3(s2)][H(s)]2

+ k−3[CH2OH(s)][CH3O(s)][s]2

Equation 5-5. Reaction rate expression of H(s)


127

d[H(s)]

dt=

1

4(−k2[H2]2[s]4 + 𝑘−2[H(s)]4)

+1

2(k−3[CH2OH(s)][CH3O(s)][s]2 − k3[HCOOCH3(s2)][H(s)]2)

− k4[CH2OH(s)][H(s)] + k−4[CH3OH][s]2 − k5[CH3O(s)][H(s)]

+ k−5[CH3OH][s]2

Equation 5-6. Reaction rate expression of CH2OH(s)

d[CH2OH(s)]

dt= k3[HCOOCH3(s2)][H(s)]2 − k−3[CH2OH(s)][CH3O(s)][s]2 + k−4[CH3OH][s]2

− k4[CH2OH(s)][H(s)]

Equation 5-7. Reaction rate expression of CH3O(s)

d[CH3O(s)]

dt= k3[HCOOCH3(s2)][H(s)]2— k−3[CH2OH(s)][CH3O(s)][s}2 + k−5[CH3OH][s]2

− k5[CH3O(s)][H(s)]

Equation 5-8. Reaction rate expression of catalytic site s

ds

dt=

1

2(−k1[HCOOCH3][s2] + k−1[HCOOCH3(s2)] + k3[HCOOCH3(s2)][H(s)]2

− k−3[CH2OH(s)][CH3O(s)][s}2) +1

4(−k2[H2]2[s]4 + 𝑘−2[H(s)]4)

+ k4[CH2OH(s)][H(s)] − k−4[CH3OH][s]2 + k5[CH3O(s)][H(s)]

− k−5[CH3OH][s]2

Equation 5-9 BDF evaluation using higher orders of Taylor polynomial

dF(t)

dt=

F(t) − F(t − h)

h

where, F(t − h) = F(t) − F′(t)h +F(2)(t)h2

2!−

F(3)(t)h3

3!+

F(4)(t)h4

4!−

F(5)(t)h5

5!

Equation 5-10. Least-squares regression function

minθE = ∑{([H2]i(t,θ)sim − [H2]

i(t)exp

)2 + ([CH3OCOH]i(t,θ)sim − [CH3OCOH]

i(t)exp

)2

Ns

i=1

+ ([CH3OH]i(t,θ)sim − [CH3OH]

i(t)exp

)2}


128

where, minθE is the minimum error; Ns is the sampling points in a batch experiment; t is the specific time of sampling, superscripts exp and sim represent experiment and simulation, respectively.

The validation results are shown in Figure 5-12, Figure 5-13, and Figure 5-14,

respectively. The absolute average relative residual (AARD, %) between the experimental

data and the model results is defined by Equation 5-11. As can be seen from Table 5-8,

the maximum AARD% for the complete set of data was 3.98%, which indicates that the

experimental results are in good agreement with the simulation model from the proposed

mechanism. In addition, the comparison between the experimental and simulation

results obtained from Figure 5-12 to Figure 5-14.

Equation 5-11. The absolute average relative residual

AARD (%) =1

N∑

|nmexp

− nmsim|

nmexp

N

1

× 100

Where, 𝑛𝑚𝑒𝑥𝑝 is the experimental moles of methanol produced, 𝑛𝑚

𝑠𝑖𝑚 is the simulation results of the moles of methanol produced, and N is the number of points.

Table 5-8. The absolute average relative residual (AARD,%) for each system

Temperature (K) Pressure (MPa) Absolute average relative residual (AARD, %) 384 2.2 3.17% 370 2.2 2.83% 346 2.2 3.98%

The regressed kinetics parameters of the proposed reaction mechanism at 346 K, 370 K

and 384 K are listed in Table 5-9. The values of 𝑘4 and 𝑘5 are very similar in each set of

experiments, implying that the reaction rates of producing methanol from the two

intermediates, CH3O(s) and CH2OH(s), are close. In addition, the reaction rate constants

increase with the increase of temperature, and this is in accordance with the observed

experimental results. The slow reaction rate observed at low temperature may also be

because hydrogen is difficult to dissolve into liquid methyl formate at low temperature

as discussed in chapter 4. The forward reaction rate of hydrogen adsorption on the active


129

sites, k2, is slower than the methyl formate adsorption step, k1, which suggests that the

slower reaction rate of the H2 adsorption than that of methyl formate. Hence, it is

assumed that the hydrogen adsorption step is the slowest step in the proposed

mechanism and we will apply the k2 values to derive the activation energy.

The robustness of the model was also checked by applying a sensitivity analysis to

determine the confidence intervals of the evaluated kinetics parameters. By introducing

± 10% disturbance on each kinetics parameter of the forward reactions, in our case, k1 to

k5, a sensitivity matrix on the ith parameter 𝑀𝜃𝑖 can be calculated by central differences

(Equation 5-12). As five kinetic parameters of forward reactions are required to test the

robustness, define 𝑁𝜃 = 5 . In each experiment, the sensitivity matrix has the scale of

𝑁𝑠 × 𝑁𝜃.

The corresponding precision matrix 𝑃 can be determined using Equation 5-13. The total

degrees of freedom 𝑁𝑑𝑓 is given by Equation 5-14. Based on the degrees of freedom, the

residual variance (𝑆𝑅2) can be determined using Equation 5-15, where 𝐸 is calculated from

Equation 5-10.

Equation 5-12. Central differences

Mθi=

∂{[H2]i(t,θ)sim +[CH3OCOH]i(t,θ)

sim +[CH3OH]i(t,θ)sim ]}

∂θi≈

[H2]i(t,θ)sim (θi

+)−[H2]i(t,θ)sim (θi

−)

2∆θi+

[CH3OCOH]i(t,θ)sim (θi

+)−[CH3OCOH]i(t,θ)sim (θi

−)

2∆θi+

[CH3OH]i(t,θ)sim ](θi

+)−[CH3OH]i(t,θ)sim ](θi

−)

2∆θi i = 1,2,3,4,5

Equation 5-13 Precision matrix P

P = (MθTMθ)

−1

Equation 5-14 Degrees of freedom

Ndf = Ns − Nθ


130

Equation 5-15 Residual variance

SR2 =

E

Ndf

Table 5-9. Regressed kinetics parameters

Temperature (K)

Forward reaction rates k1 k2 k3 k4 k5

384 275.9 ± 4.5 24.29 ± 1.70

93936 ± 6052

2086.4 ± 123.7

2649.4 ± 136.2

370 32.53 ± 1.24 9.07 ± 0.46 1331.0 ± 104.6

2722.5 ± 197.2

2061.3 ± 70.8

346 5.764 ± 0.371 2.82 ± 0.13 43.805 ± 2.187

12.84 ± 0.97

12.72 ± 0.91

Reverse reaction rates k-1 k-2 k-3 k-4 k-5

384 0.0023 1.797e5 1.434e3 1.13e-4 1.05e-4 370 1.9e-4 3.180e4 107.2 2.14e-4 2.27e-4 346 1.49e-7 8177 0.089 1.98e-3 1.72e-3

The kinetics parameters which depend on the temperature were described according to

the Arrhenius equation (Equation 5-16). The activation energy (𝐸𝑎) can be obtained from

the rate constants of the RDS at three different temperatures, which is 50.15 kJ/mol H2 in

a temperature range of 346 to 384 K and in a hydrogen pressure range of 1.8 to 2.2 MPa.

Compared with the literature data, the activation energy of 62.5 kJ/mol was concluded

from Monti et al., who applied simple power laws to determine the reaction rate in an

operating range of 1.7 to 4.5 MPa at 446 K [65]. In addition, a similar value of 53.2 kJ/mol

was summarized by Sorum at higher operating conditions with the temperature ranging

from 413 to 458 K and the pressure ranging from 3.8 to 10 MPa [156].

Equation 5-16. Arrhenius equation

ki = ko exp (−EA

RT)


131

Figure 5-12. Comparison of experimental and simulation results at T = 346 K. Operating conditions: Ptotal = 3.2 MPa, Catalyst loading = 16 g/L. Rotation speed = 800 rpm.



132


5.5.5 MECHANISM VALIDATION

It is important to verify the reaction kinetics via checking the consistency between the

simulated results and the experimental results at different pressures. In the validation

experiments, different initial pressures (1.8 MPa, 2.0 MPa and 2.2 MPa) were conducted

at 348 K, and the simulation using the obtained kinetic paraments was conducted at the

same initial conditions. The results can be found in Figure 5-15, Figure 5-16 and Figure

5-17; respectively. The open dots represent the experimental data which can be found in

Table 5-11 and the solid lines stand for the simulation results. As can be seen from Figure

5-15 to Figure 5-17, the methyl formate and hydrogen concentration decreased rapidly

with the increase in hydrogen pressures according to the Le Chatelier’s principle. It is also

found that the proposed model fits these experimental data well.


133

Figure 5-15. Validation of the modelling parameters on various pressure. Catalyst loading = 16 g/L, T = 384K. Operating conditions: PH2 = 1.8 MPa, T = 384 K, Catalyst loading = 16 g/L. Rotation speed = 800 rpm

Figure 5-16. Validation of the modelling parameters on various pressure. Operating conditions: PH2 = 2.0 MPa, T = 384 K, Catalyst loading = 16 g/L. Rotation speed = 800 rpm


134

Figure 5-17. Validation of the modelling parameter on various pressure. Operating conditions: PH2 = 2.2 MPa, T = 384 K, Catalyst loading = 16 g/L. Rotation speed = 800 rpm

Table 5-10. The absolute average relative residual (AARD%) for each system

Temperature (K) Pressure (MPa) Absolute average relative residual (AARD, %) 384 1.8 3.21% 384 2.0 2.98% 384 2.2 3.17%

Based on the experimental results and the simulation data, the absolute average relative

residual values (AARD) can be determined using Equation 5-11 and tabulated in Table 5-

10. The AARD% of this model was obtained approximately 3%. This value is reasonable

and shows that the predicted values are 3% different from the observed values. In

addition, as can be found from the Figure 5-15 to Figure 5-17, the simulation results have

good agreement on the experimental data. Based on the proposed reaction mechanism, a

schematic representation of the methyl formate hydrogenation reaction using copper

chromite is illustrated in Figure 5-18 and Figure 5-19. The step 1 and step 2 are the

adsorption steps that the H2 molecules and the methyl formate molecules are adsorbed

on the catalyst surface. As shown in Figure 5-18, hydrogen molecules adsorbed on two


135

adjacent sites of the catalysts to form H(s) and the C and H atoms from the methyl formate

attached on the active sites of the catalysts to yield CH3OCOH(s2). Subsequently, the

adsorbed species of H(s) and CH3OCOH(s2) undergo the surface reaction to produce

CH3O(s) and CH2OH(s). The produced CH3O(s) and CH2OH(s) further react with the H(s)

to produce methanol molecules as shown in Figure 5-19.

Table 5-11. Experimental values of reactants and products at different Hydrogen pressure.

1.8 MPa Time (minutes)

Moles of H2 (mole)



Conversion of H2

(%) 0 0.0384 0.172 0 0 30 ± 2.5 0.0338 0.161 0.00389 11.98 60 ± 2.5 0.0292 0.150 0.00825 23.96 90 ± 2.5 0.0264 0.143 0.0108 31.25 120 ± 2.5 0.0250 0.140 0.0121 34.90 150 ± 2.5 0.0245 0.139 0.0126 36.20 180 ± 2.5 0.0245 0.138 0.0126 36.20


Moles of H2 (mole)



Conversion of H2

(%) 0 0.0438 0.172 0 0 30 ± 2.5 0.0373 0.159 0.0058 14.84 60 ± 2.5 0.0319 0.147 0.0108 27.17 90 ± 2.5 0.0285 0.140 0.0140 34.93 120 ± 2.5 0.0269 0.138 0.0155 38.58 150 ± 2.5 0.0263 0.136 0.0161 39.95 180 ± 2.5 0.0261 0.135 0.0163 40.41


Moles of H2 (mole)



Conversion of H2

(%) 0 0.0490 0.171 0 0 30 ± 2.5 0.0393 0.153 0.00925 19.86 60 ± 2.5 0.0327 0.141 0.0155 33.27 90 ± 2.5 0.0283 0.133 0.0195 42.24 120 ± 2.5 0.0267 0.130 0.021 45.51 150 ± 2.5 0.0264 0.130 0.021 46.12


136

Figure 5-18. Schematic description of the proposed mechanism. Step 1 to Step 3

Catalyst surface

Active Sites

Catalyst surface

H-C-O-CH3

O

Catalyst surface

C O

OH

H2

Catalyst surface

H

H

STEP 1 STEP 2

H H

CH3

Catalyst surface

H2

Catalyst surface

C O

O

H

H

H

H H

CH3

STEP 3

H

Catalyst surface

C O

OH

H

H H

H

+

+-

-CH3

Adsorption Step

Surface Reaction

Catalyst surface

CH O

O

H H

H

H CH3


137

Figure 5-19. Schematic description of the proposed mechanism. Step 4 and Step 5

STEP 4 STEP 5

Catalyst surface

CH O

O

H H

H

H CH3

+

Catalyst surface

CH O

O

H H

H

H CH3

+

CH3OH

H

H

CH3OH

Catalyst surface

Catalyst surface

CH O

OH

H CH3

H H

Catalyst surface

CH OH

O

H

CH3

Chain Propagation

Chain Termination

Catalyst surface

CH O

O

H H

H

H CH3


138

5.6 CONCLUSIONS

In this chapter, the properties of copper chromite were investigated using XRD, TGA, N2

physisorption, SEM, N2O chemisorption and TPR techniques. The kinetics of methyl

formate hydrogenation reaction using copper chromite was studied at the operating

temperatures ranging from 346 to 394 K and the initial hydrogen operating pressures

ranging from 1.8 to 2.2 MPa. Both internal and external mass transfer resistance were

eliminated in all experiments. Based on the study, some conclusions can be summarised:

(1) Although trace amount of carbon monoxide and dimethyl ether were detected in

the gas phase, high selectivity of the hydrogenation reaction of 99.9 % can be

obtained.

(2) Given an identical initial total pressure of 3.2 MPa, the equilibrium conversions of

hydrogen at the temperature of 346 K, 370 K and 384 K are 83.68%, 67.45% and

46.12%, respectively. The time to reach equilibrium at temperature 346 K, 370 K

and 384 K are 140 hours, 40 hours and 2.5 hours. The solubility of hydrogen in

methyl formate decreases with the decrease of temperature, which may result in

the hydrogen is difficult to dissolve at low temperature, thus slowing down the

hydrogenation reaction of methyl formate.

(3) A reaction mechanism containing five elementary steps was proposed, and the

ordinary differential equations (ODEs) with the least-squares regression methods

were applied in MATLAB R2017b to acquire the kinetics parameters. From the

simulation results, the hydrogen adsorption rate is lower than that of methyl

formate and the hydrogen atoms desorption rate is significantly higher than that

of the adsorption, hence, we speculate that the slowest reaction step is the

hydrogen adsorption on the active species.


139

(4) Based on the kinetics parameters that were estimated from the mechanism at

different temperatures, the activation energy of the reaction was determined of

50.15 kJ/mol in the temperature range of 346 to 384 K.

(5) Furthermore, the proposed mechanism model was validated by using the

obtained parameters to compare with the experimental data at different

pressures. The simulation results and the experimental results were compared

using the average absolute relative residual (AARD %). An average of 3 % AARD

was obtained show that the simulation results have good agreement on the

experimental data, which proves that the proposed mechanism is appropriate for

the current hydrogenation reaction system using copper chromite as the catalyst.

Based on the proposed mechanism, a schematic description was provided.

Development of novel hydrogenation catalysts

140

CHAPTER 6 DEVELOPMENT OF NOVEL HYDROGENATION

CATALYSTS

6.1 INTRODUCTION

As discussed in chapter 2 and chapter 5, the hydrogenation reaction dominates the two-

step methanol synthesis process at moderate temperatures and pressures because the

reaction rate of the hydrogenation reaction is much slower than that of carbonylation

reaction at the same temperature. Therefore, investigating and developing potentially

cost-effective and efficient catalysts is crucial.

In this chapter, a design strategy of a novel catalytic system for the hydrogenation

reaction was proposed and investigated. A number of published works concluded

metallic copper provides the main active catalytic sites for methanol production via the

hydrogenation of esters [164], and thus metallic copper was incorporated in the catalytic

system. Zinc oxide (ZnO) and zirconium oxide (ZrO2) have a role of dispersing agents, and

act as addictive promoters in the catalyst system [32], [165]–[167]. Hydrotalcite-like

compound (HTC) was also selected because it acts as a carrier and a precursor to increase

the surface area of the catalytic system, which can further improve the dispersion of

active sites [9], [52], [165], [168]. Therefore, a novel catalytic system consisting of Cu,

ZnO, ZrO2 and HTC was proposed.

A simple co-precipitation method was selected to prepare this novel catalytic system, and

the characterisations and evaluation of these catalysts are discussed in this chapter. A

number of catalyst characterisation methods, including X-ray diffraction (XRD), scanning

electron microscopy (SEM), energy dispersive X-ray spectroscopy (EDX), N2

physisorption, temperature-programmed reduction (TPR), CO2-TPD, thermal


141

gravimetric analysis (TGA), X-ray photoelectron spectroscopy (XPS) and Auger electron

spectroscopy (AES), were used to investigate and compare the properties of the catalysts.

A short conclusion based on the characterisation results was made in the end.

In addition to the compound of Cu/ZnO/ZrO2/HTC catalysts, Cu/HTC, Cu/ZnO/HTC and

Cu/ZrO catalysts were designed to identify the actual function of each component in the

catalyst system for the methyl formate hydrogenation reaction. The optimised

composition of the catalytic system was studied via the hydrogenation reaction of methyl

formate. Results and some dramatic improvements were observed compared with

commercial catalysts and will be discussed in detail in Chapter 7.

6.2 DESIGN OF A NOVEL CATALYST

In Chapter 5, the kinetics mechanism of methyl formate hydrogenation reaction using a

commercial catalyst, copper chromite, was investigated. The results showed that the

hydrogen adsorption and dissociation on the catalytic sites is difficult and rate limiting.

In addition, the hydrogenation reaction rate at 370 K was very slow compared with the

carbonylation reaction at the same temperature. Thus, the slow reaction rate poses

difficulties in conducting both carbonylation and hydrogenation at the same operating

conditions. Therefore, developing new catalysts that can enhance the reaction rate at

moderate temperatures becomes important.

Based on the published works and our studies, the active catalytic species for ester

hydrogenation is copper, including Cu present in forms of Cu0, Cu+, or their combination

[95]. Hence, copper is the critical species in the novel catalytic system.

In addition to copper species, a well-dispersed copper phase may also be significant in

the catalysts to promote the reaction effectively. Therefore, some compounds which can


142

improve the copper dispersion may be necessary. ZnO is one of the most common metal

oxides that is added with the copper-based catalysts to yield the Cu/ZnO system and

increase Cu dispersion [169], [170]. It has been reported that the active Cu+ sites

dissolved in ZnO may improve the catalytic performance and thermal stability [165],

[166]. Moreover, some researchers suggested that the presence of ZnO can help increase

the copper dispersion on the surface of the catalysts as ZnO acts as a ‘physical spacer’

between copper nanoparticles [171], [172]. Therefore, in our present work, ZnO was

chosen as the copper dispersion agent in the catalyst system.

In addition, improving hydrogen adsorption on the catalysts surface is crucial in

preparing the catalytic system. As discussed in Chapter 5, the hydrogen adsorption step

on the catalytic surface is difficult, it is important to find some compounds that can

improve this adsorption step. It has been reported that ZrO2 is able to adsorb hydrogen,

thus promoting the mass transfer between gas and solid [32], [167]. This phenomenon

can be explained by hydrogen being transported from Cu or/and ZnO to ZrO2 via

hydrogen spillover effect [173], [174]. In addition, the hydrogen adsorption on the

Cu/ZnO/ZrO2 (later on, the name of Cu/Zn/Zr-HTC will be used) catalysts cannot

increase linearly with the increased amount of ZrO2, and thus the amount of ZrO2 should

be optimized when preparing catalysts [166]. Therefore, a proper ratio of ZrO2 was

selected for the novel catalysts system to improve the hydrogen adsorption on the

catalysts.

Moreover, some researchers reported that the activity of catalysts is found to be directly

proportional to the surface area, and the surface area of the catalysts can be increased by

adding some substrates with high surface area [175]. Hydrotalcite with layered brucite

structures is a potential candidate. Hydrotalcite-like compounds (HTC), also known as


143

double layered hydroxides, consist of two positive-charged brucite-like layers with

anionic compounds in the interlayer. Their general formula is [M(II)1-xM(III)x(OH)2]x+(An-

x/n)MH2O, where M(II) is bivalent metals including Mg, Zn, Co, Cu, Fe, Ni and Mn, M(III) is

trivalent metals, such as Al, Cr, Fe, and V, and A is usually CO32- ion. Due to the fact that it

has large surface area, high metal dispersion and high stability, hydrotalcite-like

compounds and their derivatives have been attracting increasing attention to be

potentially applied in chemical industries [9], [52], [165], [168]. Zhang et al. reported a

gas phase hydrogenation of ethyl acetate to produce ethanol, which was catalysed over

Ni/Al hydrotalcite-like compounds. The results showed that at 250 °C and 10 Mpa, the

selectivity and yield of the reaction were 68.2 % and 61.7 %, respectively [176].

Therefore, hydrotalcite-like compounds (HTC) can be potentially used as one major

component of preparing new catalysts to increase their surface area.

In the present work, copper, zinc, zirconium oxide and HTC were selected in the novel

catalyst system. Since HTC acts as a carrier and ZrO2 acts as a promoter, the amount of

them in the catalysts was constant in all experiments. The ratio of copper and zinc was

manipulated to optimize the catalytic performance.

6.3 CU/ZN/ZR-HTC CATALYST

6.3.1 CATALYST PREPARATION

The Cu/Zn/Zr-HTC catalysts were synthesized by a simple co-precipitation (CP) method.

The carrier material, Mg-Al layer doubled hydroxide (LDH), was pre-activated at 673 K

for 4 hours before use to enhance its surface area. A series of copper nitrates

(Cu(NO3)2·2.5H2O), zinc nitrates (Zn(NO3)2·6H2O) and zirconium nitrates

(ZrO(NO3)2·3.76H2O) with different mass ratios were dissolved in deionized water to

become a nitrate solution (0.075 g/mL). The ratio of copper nitrate/zinc


144

nitrate/zirconium nitrates in solution was controlled between 4:0:1 and 0:4:1 on a mass

basis. Precipitant, K2CO3 aqueous solution (0.38 g/mL), was prepared by dissolving K2CO3

powder in deionized water.

Typically, 1.5 g activated LDH and 10 mL nitrate solution (0.075 g/mL) were mixed under

stirring for two hours, after which 10 mL K2CO3 solution (0.38 g/mL) was added dropwise.

The obtained slurry was aged for two hours under sequential stirring. Then the pale blue

slurry was filtered and washed with deionized water by four washing steps. The obtained

precipitate was dried at 393 K in an oven overnight to yield a dried sample, followed by

grinding into a fine powder. Next, the dried fine power was calcined in air at 673 K for 4

h prior to reduction, where the colour of the catalysts became dark green. The reduction

was carried out in H2 atmosphere at 623 K for 6 hours, which allowed the copper oxides

to be reduced to metallic Cu0, and the reduced catalysts were in black colour as the size

of reduced copper was in nano-scale. The reduction temperature was determined by the

temperature-programmed reduction measurement.

The detailed information of metal compositions is given in Table 6-1. To facilitate

discussion, the groups of experiments were simply named as Cu0, Cu2, Cu4, Cu6 and Cu8.

To distinguish the catalysts at different stages, letters ‘d’ was used to represent the

catalysts after drying, ‘c’ to represent the catalysts after calcination and ‘r’ to represent

the catalysts after reduction.

Table 6-1. Metal composition of prepared catalysts

Entry Name Precursor Cu2+/Zn2+/Zr2+ ratio Cu0 HTC 0:8:2 Cu2 HTC 2:6:2 Cu4 HTC 4:4:2 Cu6 HTC 6:2:2 Cu8 HTC 8:0:2


145

6.3.2 CATALYST REDUCIBILITY

To study the reducibility or the reduction behaviour of the CuZnZr-HTC catalysts, TPR

experiments were carried out. Figure 6-1 shows the reduction profiles of all the samples

exhibiting one broad reduction peaks accompanied by shoulders from 150 to 350 °C. The

TPR profiles are deconvoluted into two Gaussian peaks. The peaks and their

corresponding contributions that derived from the deconvolution are summarised in

Table 6-2. The consumptions of H2 calculated on a mass basis for each catalyst were also

summarised in Table 6-2. It can be observed that the hydrogen consumption increases

with the increase of copper contents. This is primarily because that the CuO in the catalyst

system is reduced to active catalytic species of Cu+ and Cu0, and the amount of the active

species is proportional to hydrogen consumption.

The peak α and β may be assigned to the reduction in different aggregation states

(dispersed or bulk phase) of Cu2O, CuO and/or their combinations, representing Cu

species with different reduction capacity [177]. In the literature, the low temperature

peak (peak α) is attributed to the reduction in dispersed CuO, whereas the high

temperature peak (peak β) denotes the reduction in CuO in larger sizes [166]. Hence, it

suggests that more than one type of copper is deposited on the surface of hydrotalcite-

like compounds. As seen from Table 6-2, except cCu6, all the catalysts are dominated by

peak β at high temperatures, suggesting that great amounts of copper oxides with a large

size are present.

The TPR peaks (peak α and peak β) on the neat CuO sample appear at 231 °C and 295 °C,

which are higher than the reduction temperatures of the CuZnZr-HTC catalysts. This

indicates that there is interaction between the copper and the support (ZnO, ZrO2 and

HTC) to some extent, and such interaction leads to an easier reduction of the catalysts


146

[178]. Regarding cCu2 to cCu6, an increase amount of peak α is observed when the Zn

content decreases, while the reduction temperature of peak α on cCu6 catalysts is much

higher than those on cCu4 and cCu2. Hence, cCu6 is more difficult to reduce than cCu4

and cCu2. In addition, both peak α and peak β of cCu8 shows the lowest reduction

temperatures among other catalysts, indicating that it has the best reducibility.

Table 6-2. Centre of reduction peaks and corresponding concentrations to the TPR pattern over CuZnZr-HTC catalysts with different Cu/Zn ratio

Sample TPR peak position (°C) and concentration (%)a

H2 consumption (mmol/g)

Peak α Peak β rCu2 218 (17.1) 249 (82.9) 10.4 rCu4 210 (18.1) 253 (81.9) 17.1 rCu6 240 (63.3) 281 (36.7) 23.4 rCu8 194 (10.8) 217 (89.2) 51.7

a Values in parentheses are the contributions (%) of each peak

Figure 6-1. TPR profile of calcinated CuZnZr-HTC catalysts with different Cu/Zn ratio. (a) Cu8 (b) Cu6 (c) Cu4 (d) Cu2


147

Therefore, from the TPR profiles of the CuZnZr-HTC catalysts, both dispersed and bulk

phase CuO are present in the catalysts. The Cu8 shows the best reducibility because it has

the lowest reduction temperature among these CuZnZr-HTC catalysts. The reduction

temperature above 350 °C is required to ensure all the copper contents in the catalysts

are completed reduced.

6.3.3 CATALYST CRYSTALLINE STRUCTURE

The X-ray diffraction (XRD) measurement is used to determine the composition and

structure of the catalysts at different stages. The XRD patterns of dried CuZnZr-HTC

catalysts and the dried HTC are indicated in Figure 6-2. As can be seen, the XRD patterns

of dried HTC, or named as activated HTC, are characterised by four typically prominent

diffraction peaks (2θ), which are at 11°, 22°, 33°, 62.4° (JCPDS card #: 50-1684). Such

typical 2θ peaks of the activated HTC also remain at the CuZnZr-HTC catalysts with low

copper contents but disappear with high copper contents. This is probably due to a

replacement of Al3+ (ionic radius 0.053 nm) or Mg2+ (ionic radius 0.065 nm) by Cu2+ (ionic

radius 0.073 nm) and Zn2+ (ionic radius 0.139 nm) and the co-formation of amorphous

precipitates (such as hydroxides and hydroxyl carbonates), thus introducing distortions

on the HTC layers [179], [180].

Regarding the dried samples with high zinc contents (including catalysts Cu0 and Cu1), a

diffraction peak at 2θ = 28° can be found, which is ascribed to the crystalline Zn(OH)2

phase (JCPDS card # 38-0385). However, such a peak no longer exists when copper

contents are higher than zinc contents, revealing that either Zn(OH)2 is not generated, or

Zn(OH)2 is generated but in an amorphous form which cannot be detected from XRD.

Therefore, a large content of copper may hinder the formation of crystalline Zn(OH)2. In

addition, the crystallized Cu2(OH)3NO3 (JCPDS card # 75-1779) and Cu2CO3(OH)2 phase


148

(JCPDS card # 41-1390) are generated with the increased copper contents. The results

are consistent with the conventional co-precipitation method [181].

Figure 6-2. XRD patterns of dried CuZnZr-HTC catalyst

After the calcination at 400 °C, the CuZnZr-HTC catalysts decompose into the state of

metal oxides, which are shown in Figure 6-3. The heat treatment of the HTC at 400 °C

destroys its structure, and thus there is no characteristic peaks of activated HTCs existing

in the XRD, instead, amorphous HTC is generated which has several characteristic

diffraction peaks at 2θ = 43°, 48° and 64° [182]. In meanwhile, the sharp diffraction peaks

at 35.4° and 38.8° are detected, which are regarded as crystalline CuO [183], and the

phase crystallinity increases with the increasing copper content. In addition, there is no

peak ascribed to CuO present in both Cu0 and Cu2 sample. Moreover, amorphous ZrO2


149

exists after calcination when the ZrO2 content is low in the catalyst system, which is in

good agreement with Arena et al.’s work [184].

Figure 6-3. XRD patterns of calcinated CuZnZr-HTC catalyst

The XRD patterns of reduced CuZnZr-HTC using pure H2 at 300 °C for 6 hours are shown

in Figure 6-4. Three typical diffraction peaks at 2θ = 43.3°, 50.4° and 74.1° are detected,

which is attributed to (1 1 1), (2 0 0) and (2 2 0) crystal planes of crystalline Cu phase

(JCPDC card #04-0836). This character can be observed in Cu4 to Cu8 samples, except

Cu2. Therefore, after the hydrogen reduction, the diffraction peaks of CuO are all replaced

by those of crystalline Cu phase, indicating that the oxidized Cu species are thoroughly

reduced under hydrogen atmosphere at 300 °C. Comparing the catalyst before and after

hydrogen reduction, the intensity of crystalline ZnO remains stable, and the intensity of

diffraction peaks of ZnO increases with increased zinc content.


150

Figure 6-4. XRD patterns of reduced CuZnZr-HTC catalyst

Hence, the XRD results show that metallic copper is obtained from all the reduced

CuZnZr-HTC catalysts except rCu2. The carrier, HTC is present as its amorphous phase

after reduction. Crystalline ZnO is produced and can be detected from the XRD

measurement and ZrO2 is believed to be in its amorphous phase.

6.3.4 THERMAL STABILITY OF THE CATALYST

The thermal stability of the catalysts is determined by the thermal gravimetric analysis

(TGA). The TGA-FTIR analysis was conducted to demonstrate the thermal behaviour and

characterize the dried CuZnZr-HTC catalysts. The TGA diagrams and the DTG patterns are

shown in Figure 6-5 and the results are listed in Table 6-3. Figure 6-5 shows the tendency

of weight loss patterns of each dried CuZnZr-HTC catalyst. As the major components in

each sample are same, the characteristic weight loss at different specific temperatures is


151

similar. The percentages of total weight loss of the dried catalyst samples are also similar

as the content of metal hydroxides and carbonates in all samples were controlled at the

identical level. From the TGA, four stages can be identified regarding the thermal

decomposition of the catalysts:

1. From 30 °C to 200 °C, the weight loss is approximately 5 % and only water was

detected from FTIR. There are four major factors contributing to the weight loss

at this stage. They are the elimination of initial moisture of the samples, the

removal of weakly bound water, the decomposition of metal hydroxides, and the

removal of water in the interlayer of hydrotalcite [182], [185].

2. The major weight loss stage occurs at the temperature from 190 °C to

approximately 350 °C. The total loss of weight, ca. 17.5 % for the CuZnZr-HTC

samples mostly containing CO2 and H2O, reveals a simultaneous thermal

decomposition process of carbonate groups and hydroxyl groups from the

CuZnZr-HTC catalysts and the hydroxyl groups in the HTC. At this stage, the

deconstruction of the HTC crystal structure takes place [186]. In addition, from

the DTG patterns of Cu0 and Cu2, there are major mass losses occurring at 290 °C

due to the structure change when the ratio of zinc and copper is high. From the

TGA-DTG profile, the addition of CuO/ZnO/ZrO2 onto the HTC decreases the

decomposition temperature compared with the HTC only. This is probably due to

the fact that the introduction of metal cations into the layered structure decreases

the strength of hydrogen bonds between water molecules and interlayer anions,

thus reducing the electrostatic interaction between the layers and anions, and

leading to a lower thermal stability [187].


152

3. The weight is further lost at 500 - 600 °C. This is contributed to the destruction of

copper oxocarbonates which are formed at the second step [166].

4. The last weight loss (ca. 2-3 %) occurs at approximately 750 °C probably due to

the decomposition of CuZnZr-HTC catalysts [181]. As can be seen from Figure

6-5(a), the activated HTC does not have a major mass loss after 500 °C while all

CuZnZr-HTC catalysts have a major thermal decomposition around 850 °C, which

proves that the last mass loss is from the structure breakdown of the

CuO/ZnO/ZrO2 compounds in the catalysts. In addition, it also indicates that the

catalysts can withstand up to around 750 °C before changing their structure.

Moreover, the weight loss at the fourth step for different samples is compared in

Table 6-3. It shows that the disappearance of the weight loss in Cu0 catalysts at

this step further reveals Cu is the main contributor to this thermal decomposition.

The increased copper content results in an increase of the decomposition

temperature for the catalysts, indicating that the introduction of copper may

increase the thermal stability of the catalysts.

Table 6-3. Total mass loss of the dried catalysts

Sample Name Total mass loss Fourth step mass loss Activated HTC 39.13% 0 dCu0 27.36% 0 dCu2 28.36% 0.2% dCu4 28.95% 0.8% dCu6 27.90% 1.7% dCu8 28.83% 2.1%

In summary, the TGA patterns and the total mass loss of all the CuZnZr-HTC samples are

similar since the materials are the same. The stability of the CuZnZr-HTC catalysts

increases with the increase in the copper contents and the deconstruction temperature

of the CuZnZr-HTC catalysts is approximately 730 °C.


153

Figure 6-5. Thermogravimetry profiles of dried CuZnZr catalysts. (a) activated HTC; (b) dCu0; (c) dCu2; (d) dCu4; (e) dCu6;

(f) dCu8

6.3.5 MORPHOLOGY AND DISPERSION OF THE CATALYST

In order to discover the morphology of the catalysts as well as the dispersion, the

scanning electron microscopy with energy dispersive X-ray are used. The SEM and EDX

images of the reduced CuOZnOZrO2-HTC catalysts are shown from Figure 6-6 to Figure


154

6-10. From SEM images, it is seen that all reduced catalysts are particles placing on plate-

shaped crystals to form a layered structure with a size of 500 nm to 1 µm. The plate-

shaped crystals are characterized as the calcinated hydrotalcite compounds, which act as

the carriers of catalysts, and small particles in forms of metal or metal oxides are

deposited on calcinated hydrotalcite compounds [188].

Energy-dispersive X-ray spectrometric (EDX) analysis shows an evident segregation of

large copper agglomerates on the Cu6 and Cu8; however, such agglomerates were absent

on Cu4 and Cu2 samples. This suggests that zinc acts as a ‘spacer’ to distribute the copper

atoms well when copper is largely present [187]. Comparing Cu4 with Cu2, a better

homogenization of elements per unit of surface is observed in Cu4 samples.

Figure 6-6. SEM images and mapping of the rCu0 sample

Zn50 µm

Zr50 µm

1 µm

(a) (b)

(c)


155



Zr5 µm

Zn5 µm

Cu5 µm

1 µm

(b)

(c)

(d)

(a)

Cu5 µm

Zn5 µm

Zr5 µm

1 µm

Cu5 µm

(b)

(d)

(c)

(a)


156



The surface compositions of the catalysts, as determined by the EDX, are compared and

summarised in Table 6-4. It is clearly observed that the surface of the catalysts is enriched

in Cu and Zn at the cost of the depletion of Mg, Al and Zr. This result is in accordance with

some publications and our XRD analysis [166].

Zr10 µm

Zn10 µm

Cu10 µm

1 µm

(a) (b)

(c)

(d)

Zr25 µm

Cu25 µm

1 µm

100 nm

1 µm

(b)

(c)

(a)


157

Table 6-4. The relative surface concentration of metal (atomic %) on the CuZnZr-HTC catalysts. The values in the parentheses are the nominal concentration normalized to the total metal content of the prepared samples

Samples Cu Zn Zr Al Mg Cu0 0(0) 32.2(20.6) 2.76(3.70) 28.4(33.1) 36.6(41.9) Cu2 14.0(5.3) 28.5(15.5) 2.71(3.70) 24.2(33.1) 30.6(41.9) Cu4 22.8(10.6) 30.9(10.3) 3.04(3.70) 16.0(33.1) 27.3(41.9) Cu6 24.8(15.9) 9.90(5.17) 3.95(3.70) 27.7(33.1) 33.7(41.9) Cu8 28.5(21.2) 0(0) 3.50(3.70) 28.3(33.1) 39.5(41.9)

As can be seen from the EDX photos, all the metal compounds including Cu, Zn and Zr are

well dispersed. The SEM photos show that the Cu, ZnO and ZrO2 are successfully

deposited on the HTC surface. Based on the surface concentration of the catalysts, Al and

Mg contents are depleted while the Cu and Zn contents are enriched.

6.3.6 SURFACE AREAS AND SPECIFIC COPPER SURFACE AREA

The ASAP2010 is used to measure the specific surface area and pore volume of the

samples before and after reduction. The results are listed in Table 6-5. It reaches the

maximum value of 157.2 m2/g for cCu4, and then decreases to 141 m2/g for cCu6 but

increases only a little again to 147.1 m2/g for cCu8. After the reduction, the BET value of

the samples decreases except for cCu2. In addition, pore volume of the samples increases

with the increase of copper contents. It approaches to a maximum value of 0.37 cm3/g for

cCu4, and then decreases when the Cu/Zn ratio is higher than 1. The pore volume of all

samples increases after the reduction. The dispersion of copper (DCu) and exposed copper

surface area (SCu) were measured by reactive N2O adsorption. As can be seen from Table

6-5, the dispersion of copper is nearly the same for all groups of samples, the exposed

copper surface area, which represents the amount of active catalytic sites, increases with

increasing copper contents only expect a slight decrease in cCu6.

Table 6-5. Physicochemical properties of the calcinated samples with different Cu/Zn ratio.


158


(m2/g)

Pore volume (cm3/g)

Cu surface area b

(m2/g)

Cu dispersi

onb (%) Before

reduction After reduction a

Before reduction

After reduction a

Cu0 113.8 - 0.22 - - - Cu2 126.6 148.6 0.32 0.41 0 0 Cu4 157.2 156.0 0.37 0.52 23.30 83 Cu6 141.0 114.0 0.23 0.26 21.65 74 Cu8 147.1 144.0 0.22 0.25 47.25 80 Cu/ZrO2 23.92 26.22 0.09 0.08 14.20 23 a After reduction at 623 K. b Calculated from N2O dissociative adsorption.

In summary, compared with the CuZr catalysts without HTC, the BET surface area of the

CuZnZr-HTC catalysts is increased significantly. Hence, the large surface area can be

provided from the HTC. The Cu4 has the maximum surface area which is 156 m2/g. The

specific copper surface area of Cu2 is 0, which indicates no metallic copper is produced.

In addition, the produced metallic copper from all CuZnZr-HTC catalysts are well

dispersed on the HTC surface and the Cu8 provides the maximum copper surface area.

6.3.7 THE SURFACE BASICITY OF CUZNZR-HTC CATALYST

Figure 6-11 shows the CO2 desorption profiles of the reduced CuZnZr-HTC catalysts after

pre-treatment at 120 °C using CO2 gas over the catalysts. All profiles can be deconvoluted

into four Gaussian peaks, which are assigned to weakly (α peak), moderately (β peak) and

strongly (γ and φ peak) basic sites, respectively. The basicity of all samples is listed in

Table 6-6.

Regarding the catalysts after reduction, the weakly basic sites and moderate basic sites

are related to surface hydroxyl group (OH-) and the metal-oxygen pairs (e.g. Zn-O, Al-O,

and Zr-O etc.), respectively [189], [190]. As can be seen from Table 6-6 and Figure 6-11,

with increasing copper contents, peak α and β are shifted from 167 to 183 °C and 213 to

255 °C, respectively. This is due to the interaction of ZnO and ZrO2, which is in agreement


159

with Huang et al.’s work [188]. It is known that the surface of ZrO2 has Lewis basic sites

to adsorb CO2. It is expected that the alkaline ZnO enhances the affinity of the system to

CO2 [171], [172], [177]. Hence, the total number of weakly basic sites decease when the

Cu/Zn ratio increases. The two strongly basic sites (γ and φ peak) are attributed to low-

coordination oxygen atoms [173], [189]. The γ peak occurred at high temperatures from

590 °C to 600 °C is ascribed to the Mg-Al structure existing in the HTC precursor, which

has been proved by Gao et al [168].

The total number of basic sites per catalyst on the reduced CuZnZr-HTC catalysts tend to

increase up to 469 µmol/g for rCu4 and decrease down to 367 µmol/g for rCu6, but

increase again to 496 µmol/g for rCu8. This is probably because the interaction among

components is changed depending on the compositions, which further influences the

electronic effects and have significant impacts on the surface basic sites [177].

Table 6-6. The amount and distribution of basic sites of the reduced CuZnZr-HTC catalysts

Sample CO2-TPD peaks position (°C) and concentration (%)a Number of basic sites (µmol/g)

Peak α Peak β Peak γ Peak φ

rCu2 167 (24.7) 213 (40.9) 343 (27.1) 637 (7.3) 418 rCu4 169 (23.4) 237 (34.3) 435 (24.3) 592 (18.1) 469 rCu6 182 (18.1) 236 (39.5) 404 (33.2) 633 (9.3) 367 rCu8 183 (18.7) 255 (37.9) 451 (43.4) - 496

a Values in parentheses are the contributions (%) of each peak


160

Figure 6-11. CO2-TPD pattern of the reduced CuZnZr-HTC catalysts. (a) Cu8; (b) Cu6; (c) Cu4; (d) Cu2

To conclude, the basic sites of the CuZnZr-HTC catalysts are very close and the Cu8

present the maximum number of basic sites of 469 µmol/g. Some researches believed

there is a correlation between the catalysts basic sites and the products selectivity, which

will be discussed in chapter 7.

6.3.8 CHEMICAL STATES OF ELEMENTS IN THE CATALYST

The XPS measurement of the catalysts was conducted to determine the chemical states of

the elements and the XPS spectra of Cu, Zn, Zr and O for series of CuZnZr-HTC catalysts


161

are presented in Figure 6-12 to Figure 6-15. The binding energies (BE) and the full width

at half maximum (FWMH) are summaries in Table 6-7 to Table 6-9.

It can be seen from Table 6-7, these calcinated samples exhibit Cu 2p3/2 and Cu 2p1/2 main

peaks in the BE ranges from 933.3 to 933.9 eV and from 963 to 953.7 eV with a spin-orbit

coupling energy gap of 20 eV, respectively. From Figure 6-12, it is also noticed that both

Cu 2p3/2 and Cu 2p1/2 peaks are accompanied by intense satellite (Cu2+) featuring at 942

eV and 962 eV for all calcinated samples, and the Cu2+ satellite disappears after reduction,

which proves that the copper oxide is reduced to either Cu+ or a metallic state of Cu. Since

the BE value of Cu0 (932.6 eV) and Cu+ (932.4 eV) are close, it is difficult to distinguish by

typical XPS. Hence, the Auger electron spectrum of L3M45M45 (hereafter referred as Cu

LMM) is used to determine the states of copper. Detailed discussion of the distribution of

Cu+ and Cu0 will be given later.

As can be seen from Table 6-8 and Figure 6-13, the Zn 2p3/2 peak exhibits at the BE values

of between 1021.8 and1022.1 eV for both calcinated and reduced CuZnZr-HTC samples

with a FWHM value of 1.84 ± 0.08 eV. These characteristic values are close to a reference

of ZnO cat the BE values of 1021.3 to 1022.0 eV and a FWHM value of 2.0 eV [191].

Therefore, it is speculated that the chemical environment of Zn is not affected by other

components of the catalysts. In addition, since the chemical state of the Zn is unchanged

before and after reduction, indicating that ZnO is stable in the catalysts, which is

consistent with the XRD results where there are ZnO crystals present before and after

reduction.

From Zr 3d core level XP spectra of CuZnZr-HTC samples (shown in Figure 6-14) and the

XPS parameters summarized in Table 6-9, Zr exhibits the spin-orbit doublet of the 3d core

level into 3d5/2(182.04 – 182.3 eV) and 3d3/2 (184.37 – 184.67 eV) levels with an energy


162

gap of 2.37 ± 0.02 eV between them. These parameters are in good agreements with the

XP spectra of ZrO2, indicating that Zr4+ is present in the catalysts, and the state of ZrO2

remains unchanged before and after reduction [192].

Table 6-7. XPS parameters of Cu core level in CuZnZr-HTC catalysts

Samples Cu 2p3/2 Cu 2p1/2 BE (eV) FWHM (eV) BE (eV) FWHM (eV)

cCu2 933.4 2.73 953.0 2.88 cCu4 933.7 3.03 953.7 3.33 cCu6 934.0 2.89 953.7 2.98 cCu0 934.1 3.12 953.9 3.09

rCu2 933.0 2.11 952.9 2.79 rCu4 932.8 1.58 952.7 2.10 rCu6 932.8 1.51 952.6 2.12 rCu8 932.8 1.29 952.6 1.77

Table 6-8. XPS parameters of Zn core level in CuZnZr-HTC catalysts

Samples Zn 2p3/2 Zn 2p1/2 BE (eV) FWHM (eV) BE (eV) FWHM (eV)

cCu2 1021.9 1.95 1045.1 1.92 cCu4 1021.8 1.89 1044.9 2.03 cCu6 1021.9 1.87 1044.9 2.11

rCu2 1022.0 1.78 1045.2 1.99 rCu4 1022.0 1.72 1045.1 1.92 rCu6 1021.9 1.85 1045.0 2.03

Table 6-9. XPS parameters of Zr core level in CuZnZr-HTC catalysts

Samples Zr 3d5/2 Zr 3d3/2 BE (eV) FWHM (eV) BE (eV) FWHM (eV)

cCu2 182.1 1.44 184.5 1.42 cCu4 182.0 1.40 184.4 1.37 cCu6 182.3 1.34 184.6 1.46 cCu8 182.3 1.43 184.6 1.48

rCu2 182.2 1.34 184.6 1.42 rCu4 182.3 1.29 184.7 1.32 rCu6 182.1 1.29 184.5 1.32 rCu8 182.4 1.36 184.8 1.39

The O 1s broad peaks can be decomposed in two peaks at the corresponding positions

using XPS peak splitting program (XPS Casa Software), whose relative contents are

shown in Figure 6-15. It is shown that the O 1s peaks undergo a significant change after


163

reduction. Based on a number of studies [193]–[195], there may be two types of oxygen

in the catalyst system after reduction: oxygen species existing in ZrO2, Cu2O and ZnO and

oxygen species existing in Zr-OH. The binding energy for the two types of oxygen is in a

range of 529.8 to 530.6 eV and a range of 531.7 to 532.2 eV, respectively. It is shown that

the oxygen species with BE value higher than 531 eV are attributed to OH group,

chemisorbed oxygen or carbonates groups [195], [196]. Hence, the oxygen species tends

to transform into O2- species after reduction.

To distinguish Cu0 from Cu+ species in a better way, the X-ray Induced Auger electron

spectra of calcinated and reduced samples in a kinetic region of 906 eV to 920 eV are

presented in Figure 6-16. The position of peaks can be distinguished as Auger Cu LMM

transit Cu2+ in a kinetic energy range of 913 to919 eV [197]. The shifted position of the

Cu Auger has agreements with the results previously reported in the literature, where

such position shifts are attributed to the chemical position and the specific bonding

interactions between the oxide phases species [195], [197]. In our work, peaks at c.a. 918

eV and c.a. 913 eV are detected, which are assigned to Cu0 and Cu+, respectively. As can

been seen from the Table 6-10, the main peaks for cCu2, at a position of 914.2 eV is shifted

to a higher kinetic energy value up to 918.4 eV for cCu8, with increasing Cu contents [195].

This is in accordance with the highest Cu content of the EDX results over cCu8 samples.

A new parameter named modified Auger parameter (αCu) is introduced to determine the

chemical state of copper in the samples [195], [197]. This parameter is defined by

Equation 6-1.

Equation 6-1. Auger parameter (αCu)

αCu = EB + EK


164

where, EB is the binding energy of Cu 2p2/3 (2p core level) and Ek is the kinetic energy of the Cu LMM Auger electron.

The αCu values are at approximately 1847.6 and 1852 eV (Table 6-10) corresponding to

Cu+ and Cu2+ species, respectively. The α value of 1847.6 eV for cCu2 means that only Cu+

is present, while two α values associated with other catalysts imply two types of Cu

species are present, but Cu2+ is still a dominant species in all catalysts. Those results are

consistent with the XRD patterns.

After hydrogen reduction, the Cu species changed their chemical states. In Table 6-11, the

αCu values of reduced samples are at c.a. 1847 and 1850 eV, which corresponds to Cu+ and

Cu0, respectively. Only Cu+ are present in rCu2, while both types of Cu0 and Cu+ are

observed in other reduced groups. The value of Cu0/Cu+ for all other reduced groups is

less than 1, indicating that Cu+ is still the major species. The results are consistent with

the XRD results of both calcined and reduced results. In addition, only Cu+ phase in the

rCu2 sample spears after reduction, whereas other reduced groups contain Cu0 phase,

indicating that the reduction of CuO is easier than that of Cu2O, which is consistent with

Kim et al.’s work [198].

Table 6-10. XPS parameters of calcined CuZnZr-HTC samples

Samples Cu 2p3/2 (eV) Cu Auger (eV) αCu (eV) Cu+/ Cu2+

cCu2 933.4 914.24 1847.6 1 (all Cu+) cCu4 933.7 913.94/918.08 1847.64/1851.78 2.76a

cCu6 933.9 913.58/918.3 1847.48/1852.20 1.60 cCu8 934.0 914.70/918.4 1848.70/1852.40 1.83

a Defined by the ratio of the peak area

Table 6-11. XPS parameters of reduced CuZnZr-HTC samples

Samples Cu 2p3/2 (eV) Cu Auger (eV) αCu (eV) Cu0/Cu+ rCu2 933.0 913.96 1847.0 0 (all Cu+) rCu4 932.8 914.41/917.8 1847.21/1850.60 0.64 rCu6 932.8 913.27/917.31 1846.07/1850.11 0.63 rCu8 932.8 914.21/917.3 1847.01/1850.10 0.66


165

In addition, it is noted that a new peak appears at a low kinetic energy position (c.a. 910

to 910.3 eV) for all the samples except rCu6 and rCu8. Such peak was also detected in

Zhan et al.’s work [183], which is attributed to Cuα+ in a perovskite structure (A2+B4+O3, A

and B are metals). Hence, it is likely to form a perovskite structure in the prepared

catalysts, which is mostly in the presence of CuZrO3 compound. However, based on our

XRD results, there is no indication that such crystalline structure is present in the samples,

probably because the perovskite structure is not crystalline as the amount of zirconium

is limited. Moreover, it is found such peak disappears when copper contents increase,

similar phenomenon was found in Saha et al.’s work, where it was observed that large

stoichiometric ratio of copper in CuZrO3 resulted in structure defects [199].

In short, the XPS spectrum show that ZnO and ZrO2 in the CuZnZr-HTC catalysts stay

unchanged before and after reduction. From Auger spectrum, no Cu0 phase is detected

from the Cu2 catalysts, which is consistent with the XRD, N2O chemisorption results. In

addition, the ratio of Cu0/Cu+ in the Cu4 to Cu8 catalysts is very similar.


166

Figure 6-12. Cu 2p core level X-ray photoelectron spectra of CnZuZr-HTC series samples. (A) Cu2; (B) Cu4; (C) Cu6; (D) Cu8 (i) represents calcinated state, (ii) represents reduced state.

NameCu 2p3/2Cu 2p1/2

Pos.932.753952.591

%Area49.1050.90

Cu

2p

x 103

15

20

25

30

35

40

CP

S

970 965 960 955 950 945 940 935 930 925Binding Energy (eV)


Pos.932.813952.649

%Area49.8950.11

Cu

2p

x 103

20

25

30

35

40

45

CP

S


rCuZnZr(2:6:2)-HTC


Pos.933

952.861

%Area47.9752.03

Cu

2p

x 102

140

150

160

170

180

190

200

210

CP

S



Pos.933.703953.662

%Area55.3044.70

Cu

2p

Cu

2p

x 103

18

20

22

24

26

28

30

CP

S



Pos.933.974953.728

%Area54.4945.51

Cu

2p

Cu

2p

x 103

12

14

16

18

20

22

24

26

CP

S


cCuZnZr(2:6:2)-HTC

Name

Cu 2p3/2

Cu 2p1/2

Pos.

933.374

953.061

%Area

50.02

49.98

Cu

2p

x 102

150

160

170

180

190

200

CP

S



Pos.934.073953.882

%Area53.3946.61

Cu

2p

Cu

2p

x 103

14

16

18

20

22

24

26

28

30

32

CP

S



Pos.932.792952.591

%Area50.2249.78

Cu

2p

x 103

15

20

25

30

35

40

45

50

55

60

CP

S


(C-i)

(B-i)

(A-i)

(D-i)

(C-ii)

(A-ii)

(B-ii)

(D-ii)


167

Figure 6-13. Zn 2p core level X-ray photoelectron spectra of CnZuZr-HTC series samples. (A) Cu2; (B) Cu4; (C) Cu6; (D) Cu8 (i) represents calcinated state, (ii) represents reduced state.

NameZn 2p3/2Zn 2p1/2

Pos.1021.851044.94

%Area49.3150.69

Zn

2p

x 103

20

22

24

26

28

30

32

34

36

38

CP

S

1050 1045 1040 1035 1030 1025 1020 1015Binding Energy (eV)


Pos.1021.891044.97

%Area49.0950.91

Zn

2p

x 103

20

22

24

26

28

30

32

34

36

38

40

CP

S


rCuZnZr(2:6:2)-HTC


Pos.1022.071045.19

%Area49.8850.12

Zn

2p

x 103

20

30

40

50

60

70

80

90

CP

S


cCuZnZr(2:6:2)-HTC


Pos.1021.881045.12

%Area59.1840.82

Zn

2p

Zn

2p

x 103

20

30

40

50

60

70

80

90

CP

S



Pos.1021.811044.89

%Area50.0349.97

Zn

2p

x 103

30

40

50

60

70

80

90

CP

S



Pos.1022.011045.09

%Area49.6750.33

Zn

2p

x 103

30

40

50

60

70

80

90

100

CP

S


(C-i)

(B-i)

(A-i)

(C-ii)

(A-ii)

(B-ii)


168

Figure 6-14. Zr 3d core level X-ray photoelectron spectra of CnZuZr-HTC series samples. (A) Cu2; (B) Cu4; (C) Cu6; (D) Cu8 (i) represents calcinated state, (ii) represents reduced state.

NameZr 3d5/2Zr 3d3/2

Pos.182.039184.373

%Area50.5649.44

Zr

3d

x 102

22

24

26

28

30

32

34

36

38

CP

S

189 186 183 180 177Binding Energy (eV)


Pos.182.304184.677

%Area49.9650.04

Zr

3d

x 102

25

30

35

40

45

50

55C

PS


rCuZnZr(2:6:2)-HTC


Pos.182.202184.582

%Area49.5550.45

Zr

3d

x 102

20

25

30

35

40

45

CP

S


cCuZnZr(2:6:2)-HTC


Pos.182.126184.462

%Area51.3448.66

Zr

3d

x 102

22

24

26

28

30

32

34

36

38

CP

S



Pos.182.131184.523

%Area48.3851.62

Zr

3d

x 102

15

20

25

30

35

40

45

50

55

CP

S



Pos.182.254184.638

%Area49.9850.02 Z

r 3

d

x 102

20

30

40

50

60

70

80

CP

S



Pos.182.281184.603

%Area50.9349.07

Zr

3d

x 102

20

25

30

35

40

45

50

55

CP

S



Pos.182.447184.818

%Area49.9350.07

Zr

3d

x 102

20

30

40

50

60

70

80

90

CP

S


(C-i)

(B-i)

(A-i)

(D-i)

(C-ii)

(A-ii)

(B-ii)

(D-ii)


169

Figure 6-15. O 1s core level X-ray photoelectron spectra of CnZuZr-HTC series samples. (A) Cu2; (B) Cu4; (C) Cu6; (D) Cu8 (i) represents calcinated state, (ii) represents reduced state.

cCuZnZr(2:6:2)-HTC

NameO2-OH/CO

Pos.530.636

532.24

%Area52.6647.34

O 1

s

x 102

90

100

110

120

130

140

150

160

170

180

CP

S

537 534 531 528Binding Energy (eV)

rCuZnZr(2:6:2)-HTC

NameO2-OH/CO

Pos.530.629532.174

%Area65.0234.98

O 1

s

x 103

10

12

14

16

18

20

22

24

26

28

30

CP

S


NameO2-OH/CO

Pos.530.133531.783

%Area32.7767.23

O 1

s

x 103

8

10

12

14

16

18

20

22

24

26

28

CP

S


NameO2-OH/CO

Pos.530.573

532.09

%Area60.8239.18 O

1s

x 103

10

15

20

25

30

35

CP

S


NameO2-OH/CO

Pos.529.88531.74

%Area37.4562.55

O 1

s

x 103

6

8

10

12

14

16

18

20

22

24

CP

S


NameO2-OH/CO

Pos.530.252531.773

%Area52.1947.81

O 1

s

x 103

10

15

20

25

CP

S


NameO2-OH/CO

Pos.529.942531.703

%Area42.1457.86

O 1

s

x 103

5

10

15

20

CP

S


NameO2-OH/CO

Pos.530.494532.047

%Area59.4540.55

O 1

s

x 103

4

6

8

10

12

14

16

18

20

22

CP

S


(C-ii)

(A-ii)

(B-ii)

(D-ii)

(C-i)

(B-i)

(A-i)

(D-i)


170

Figure 6-16. X-ray induced Auger electron spectra of catalysts. (a) cCu2; (b) rCu2; (c) cCu4; (d) rCu4; (e) cCu6; (f) rCu6; (g) cCu8; (h) rCu8

x 102

95

100

105

110

115

120

125

130

135

140

Inte

nsi

ty

908 912 916 920Kinetic Energy (eV)

x 102

90

95

100

105

110

115

120

125

Inte

nsi

ty908 912 916 920

Kinetic Energy (eV)

x 102

100

105

110

115

120

125

130

135

Inte

nsi

ty

908 912 916 920 924Kinetic Energy (eV)

x 102

105

110

115

120

125

130

135

140

145

150

Inte

nsi

ty


x 102

65

70

75

80

85

90

Inte

nsi

ty


x 102

70

75

80

85

90

95

Inte

nsi

ty


(a) (b)

(e) (f)

(c) (d)

x 102

60

70

80

90

100

110

Inte

nsi

ty

904 908 912 916 920 924Kinetic Energy (eV)

x 102

60

70

80

90

100

Inte

nsi

ty

904 908 912 916 920 924Kinetic Energy (eV)

(g) (h)


171

6.4 CONCLUSIONS

The CuZnZr-HTC catalysts with a series of mass ratio of Cu2+:Zn2+:Zr4+ with x:y:2 (x+y =

8, x=0,2,4,6,8) were successfully prepared by a simple co-precipitation method. The

physical and chemical properties of the catalysts with different ratio of copper/zinc were

investigated by XRD, N2 physisorption, SEM, N2O chemisorption, TPR, CO2-TPD and XPS

techniques. Some conclusions can be summarised:

(1) The Cu/Zn/Zr-HTC catalysts were synthesised and can be thoroughly reduced at

pure hydrogen atmosphere above 623 K using the temperature-programmed

reduction (TPR) measurement, therefore, the reduction temperature of the

catalysts was determined at 623 K.

(2) Both metallic copper (Cu0) and Cu+ can be detected from sample rCu4, rCu6 and

rCu8 and only Cu+ species was detected in rCu2 sample from XRD analysis and

Auger electron spectrum(AES). The AES results also indicate that the ratio of

Cu0/Cu+ of rCu4, rCu6 and rCu8 samples are very close. The actual catalytic sites

for hydrogenation reaction will be discussed in chapter 7.

(3) Crystalline ZnO was detected from sample rCu2, rCu4 and rCu6 by XRD analysis.

The chemical state of zinc element analysis from the XPS measurement shows that

the ZnO phase remains unchanged before and after reduction.

(4) ZrO2 is present in the amorphous phase as its crystalline phase cannot be detected

from the XRD analysis, while the XPS analysis shows the existence of ZrO2.

(5) The rCu4 provides the largest BET surface area (156 m2/g) among the catalysts

from the N2 adsorption/desorption process. However, the rCu8 presents the

largest specific copper surface area, which is 47.25 m2/g from the N2O dissociative

measurement.

Determination of the catalytic performance of the novel hydrogenation catalysts

172

CHAPTER 7 DETERMINATION OF THE CATALYTIC

PERFORMANCE OF THE NOVEL HYDROGENATION

CATALYSTS

7.1 INTRODUCTION

As discussed in Chapter 6, a catalytic system containing copper, zinc and zirconium

deposited on hydrotalcite-like compounds was designed for the hydrogenation reaction

of methyl formate. Several characteristics of the catalytic system, including thermal

stability, reducibility, structure, surface dispersion, number of basic sites, surface area,

copper surface area, and elemental states of the catalysts, were studied intensively.

In this chapter, the roles of each component in the catalyst system were investigated via

different catalyst systems including ZnO/ZrO2-HTC, Cu-HTC, Cu/ZnO-HTC, Cu/ZrO2-HTC

and Cu/ZrO2 without HTC for the hydrogenation reaction at 384 K. The catalytic effect of

the ratio of Cu/ZnO contained in the CuZnZr-HTC catalytic system was studied on the

methyl formate hydrogenation reaction at 384 K. The reaction rates were compared

using the pseudo space time yield (STYPS), and the product selectivity was also

determined.

The results show that the bimetallic catalysts coated with hydrotalcite-like compound

(Cu/ZrO2-HTC) has the best catalytic performance at 384 K. A significant improvement

on the reaction rate was achieved using Cu/ZrO2-HTC at a lower temperature (370 K)

compared with the commercial catalyst, copper chromite, used for ester hydrogenation

reactions. A discussion on the advantages of using the Cu/ZrO2-HTC is given at the end of

the chapter.


173


The experimental apparatus and procedure are the same as described in section 5.2 and

the experiments conducted in this chapter were listed in Table 7-1.

Table 7-1. Experiment Operating conditions

Parameters Experimental conditions Catalysts Cu/ZnO/ZrO2-HTC system Rotation speeds (rpm) 800 Catalyst loadings (g/L) 16 Temperatures (K) 370 and 384 Hydrogen pressure (MPa) 2.2

7.3 RESULTS AND DISCUSSION

7.3.1 ROLES OF THE COMPONENTS IN THE CU/ZNO/ZRO2-HTC CATALYTIC SYSTEM

The roles of each in the catalytic system are important and can be determined by

comparing the reaction rate using different combinations of the components. Hence, five

different catalysts including ZnO/ZrO2-HTC, Cu-HTC, CuO/ZnO-HTC, Cu/ZrO2-HTC and

Cu/ZrO2 without HTC were selected and their compositions were summarised in Table

7-2. The total pressure versus the reaction time of each experiment was recorded and

shown in Figure 7-1. For easy interpolation, a reaction time span of 200 mins was

selected.

Table 7-2. Metal compositions of prepared catalysts

Entry Name Precursor Cu2+/Zn2+/Zr2+ ratio ZnO/ZrO2-HTC HTC 0:8:2 Cu-HTC HTC 8:0:0 Cu/ZnO-HTC HTC 8:2:0 Cu/ZrO2-HTC HTC 8:0:2 Cu/ZrO2 only N.A. 8:2


174

Figure 7-1. The pressure profile of the catalysts. The total pressure of the system, including the partial pressure of the solvent and the pressure of the gas.

Role of copper component

As can be found from Figure 7-1, the catalysts without copper species do not have

catalytic performance compared to other prepared catalysts containing Cu species have

catalytic performance on the hydrogenation reaction of methyl formate. Therefore, the

copper species in the catalysts provide the active sites for the hydrogenation reaction.

Role of hydrotalcite

A slow pressure drop from the Cu/ZrO2 catalysts is found, compared with the other

copper-based HTC catalysts. The low reaction rate of the Cu/ZrO2 only catalyst might be

attributed to its small surface area. As discussed in section 6.4, the surface area of the

Cu/ZnO/ZrO2-HTC after reduction can attain 145 m2/g, however, the Cu/ZrO2 catalyst


175

only has a BET surface area of 20.7 m2/g after reduction. This implies that HTC serves a

role of increasing the surface area and further increasing the reaction rate.

In addition, some metal copper particles which have inactive species and poor stability

can be observed after the reaction using Cu/ZrO2, which is shown in Figure 7-2. This

phenomenon illustrates that the large size of Cu synthesized from the co-precipitation

method is not dispersed well without the presence of HTC. Hence, HTC may also improve

the dispersion of copper oxide particles during the preparation stage of the catalysts.

Figure 7-2. The appearance of the Cu/ZrO2 catalysts after reaction

Role of ZnO and ZrO2

The Cu-HTC catalysts show the slowest pressure decline rate, compared with the

Cu/ZrO2-HTC and Cu/ZnO-HTC, indicating that the addition of ZnO and/or ZrO2

components is advantageous. The advantages, such as the enhanced dispersion of the

copper/copper oxides [165], [166], the spillover effect on the copper from ZrO2 [171],

[172], the improved thermal stability [32], [167], have been found in a number of

published works.

1 cm


176

In a comparison between the Cu/ZrO2-HTC and the Cu/ZnO-HTC catalysts, a better

performance was obtained using the Cu/ZrO2-HTC catalysts. This suggests that the

function of ZrO2 may outweigh ZnO when HTC is present. This can be explained by the

‘spillover’ effect from the modified ZrO2 in which ZrO2 may improve the mass transfer

between the hydrogen molecules and the surface of the catalysts, thus further increasing

the reaction rate. In addition, the dispersion role provided by the ZnO may be

compromised by the HTC to some extent as the HTC may be able to provide large surface

area and also improves the copper dispersion.

7.3.2 BY PRODUCTS FORMATION

As indicated from section 7.3.1, no pressure decline can be found using the Zn/Zr-HTC

catalysts, implying that there is no reaction occurring. The results are further confirmed

through the GC analysis as no methanol was detected in both liquid samples and gas

samples.

For the gas products produced from the hydrogenation reaction using the Cu/ZnO/ZrO2-

HTC catalysts, only a trace amount of dimethyl ether and carbon monoxide were detected

by GC, which is similar to the commercialized catalyst, copper chromite. As explained in

Chapter 5, two possible side reactions may occur, including a reversed decarbonylation

reaction (shown in Reaction 7-1) where CO is produced and a dehydration reaction of

methanol (Reaction 7-2) where dimethyl ether and water are generated. Since the

amounts of by-products are small and can be neglected, the selectivity of the

hydrogenation reaction using the Cu/ZnO/ZrO2-HTC catalysts is 99.9%.

Reaction 7-1. Decarbonylation reaction of methyl formate

CH3OCOH ⇌ CH3OH + CO ∆H = 16.8 kJ/mol

Reaction 7-2. Dehydration of methanol


177

2CH3OH ⇌ H2O + CH3OCH3 ∆H = 51.3 kJ/mol

In addition, some published work correlates the basic sites of catalysts with the

selectivity of products while they stated that the more basic sites of the catalysts, the

higher selectivity of the desired products. However, there is no evidence in our study

showing a correlation between these two factors.

7.3.3 THE CATALYTIC EFFECT OF THE RATIO OF CU/ZNO ON THE REACTION

EFFECTS ON THE REACTION RATE AND THE TIME TO REACH EQUILIBRIUM

The liquid and gas products using the catalysts of rCu2, rCu4, rCu6 and rCu8 were

analysed via the GC, and the compositions of rCu2, rCu4, rCu6 and rCu8 were given in

Table 6-1. The concentration profile of the product methanol and the reactant hydrogen

over the reaction time are shown in Figure 7-3 and Figure 7-4, respectively. Since the

experimental system is a batch reactor, equilibrium is finally attained at ‘long’ times.

Therefore, the overall time to achieve the equilibrium is used to determine the reaction

rate, which are listed in Table 7-3. In addition, in order to compare the catalytic

performance, a new term called the pseudo-space time yield (STYPS), which is defined as

the amount of methanol produced per gram of catalysts in the first 100 minutes, is given

in Equation 7-1. The STYPS of each catalyst is also provided in Table 7-3.

From Figure 7-3 and Figure 7-4, it can be found that using the rCu2 catalyst takes the

longest time to reach the equilibrium. Even after 600 minutes the reaction only reaches

half of its equilibrium conversion. With increasing the copper contents in the catalysts,

the time to achieve the equilibrium is reduced significantly, and the reaction is

accelerated correspondingly. It is observed that the reaction using the rCu4 and the rCu6

has a similar rate, but the reaction is accelerated dramatically when using the rCu8

catalyst. From STYPS analysis, consistent results can also be found. The real-time reaction


178

rate at the first 100 minutes using the rCu2 is the slowest, and the rCu4 and the rCu6

achieve a comparable reaction rate with only a slightly slower rate when using rCu6;

however, the reaction rate using the rCu8 increases dramatically to 4.3 g methanol/g

catalyst compared to other catalysts in the first 100 minutes. It is speculated that the

surface area of active metallic copper and the interaction between ZrO2 and Cu have

direct effects on the hydrogenation reaction.

Equation 7-1. Pseudo-space time yield (STYPS)

STYPS =methanol (g)

catalyst amount (g) in the first 100 minutres

Table 7-3. Catalytic performance for hydrogenation of methyl formate

Catalysts Time to equilibrium STYPS for first 100 minutes (g methanol/g catalyst)

rCu2 1080 minutes 0.82 rCu4 750 minutes 2.25 rCu6 720 minutes 1.86 rCu8 150 minutes 4.3


179

Figure 7-3. The amount of methanol produced over time

Figure 7-4. The amount of hydrogen in the reactor over time

THE EFFECTS OF BET SURFACE AREA ON CATALYTIC PERFORMANCE

Compared with the surface area of the copper-based HTC catalysts after reduction (Table

6-5), they are very close, but the reaction rates are different. Hence, a direct comparison

between the BET surface area of catalysts and the catalytic performance may not be

sufficient. In addition of the BET surface area, the amount and the dispersion of active

metallic copper may be also important to evaluate the performance of the catalysts.

CATALYTIC PERFORMANCE RELATED TO SPECIFIC COPPER SURFACE AREA AND

COPPER DISPERSION

As discussed in Chapter 6, the copper surface area is an important factor in the copper-

based catalytic reactions [169], [175]. A relationship of the copper surface area and the

STYPS for rCu2, rCu4, rCu6 and rCu8 is shown in Figure 7-5. The STYPS increases with the

increase of the exposed copper surface area for all hydrogenation reactions at 384 K, and


180

a nearly linear relationship can be found. This is probably because a large value of SCu

provides more active sites, which in turn have more opportunities to adsorb and store H2.

The adsorbed H2 is further dissociated in the form of atomic hydrogen, and transported

from Cu and/or ZnO to ZrO2 via ‘spillover’ effect [32], [83]. It is noted that rCu8 provides

the maximum copper surface area and STYPS, however there is no ZnO contained in the

catalytic system, which indirectly indicates that the Cu has more profound ‘spillover’

effects than that of ZnO. Therefore, to improve the catalytic activity for the hydrogenation

reaction of methyl formate, the copper-zirconium catalysts obtained using the current

simple co-precipitation method may be optimized by changing the ratio of copper and

zirconium in the future work.

In addition, it is also found that the prepared catalysts using the simple co-precipitation

method are stable and the random error is minimal, since the N2O dissociative analysis

shows that the dispersion of metallic copper among the rCu4, rCu6 and rCu8 is very close

when the designed mass ratio of Cu/Zn is above 1 (rCu2 does not contain metallic copper).


181

Figure 7-5. The relationship between the space time yield and copper surface area

CATALYTIC EFFECTS OF THE ACTIVE SITES CU0 AND CU+

Figure 7-3 and Figure 7-4 show that the rCu8 catalysts have the best catalytic

performance compared with other groups of catalysts. This can be attributed to the large

specific copper surface area and good copper dispersion. However, as indicated in the

XPS and XRD study, the rCu2 catalysts only contain the Cu+ species rather than the active

metallic copper, but it can also promote the hydrogenation reaction implying that the Cu+

species may also be the active catalytic sites for the hydrogenation reaction. In addition,

some researches outlined that the ratio of Cu0/Cu+ is another important factor that may

affect the reaction rate [195]. As shown in the XPS analysis, the ratio of Cu0/Cu+ are very

close (approximately 0.66) for the rCu4, rCu6 and rCu8 catalysts, but the reaction rate of

each catalyst system is not the same as expected, indicating that it may not be sufficient

to show the relationship between the Cu0/Cu+ and the reaction rate in the hydrogenation


182

reaction. Hence, both Cu+ and Cu0 are the active sites in the current catalyst system to

accelerate the hydrogenation reaction.

7.4 Comparison of Cu/ZrO2-HTC catalysts with copper chromite

To evaluate the commercial potential of the developed catalyst system, it is important to

compare the best catalysts from the Cu/ZnO/ZrO2-HTC catalytic system with copper

chromite, which is one of the most common catalysts used for the ester hydrogenation

reactions. A comparison including the characteristics of catalysts and the catalytic

performance between rCu8 and copper chromite was conducted for the hydrogenation

reactions of methyl formate at 370 K and 384 K.

7.4.1 THE CHARACTERISTICS OF CATALYSTS

7.4.1.1 The surface area and dispersion of active metallic copper

Table 7-4 summarises the physicochemical properties of copper chromite and rCu8 (or

named CuZr(8:2)-HTC) catalysts. It is found that the BET surface area and the Cu surface

area are both very similar; however, the copper dispersion of copper chromite is less than

the rCu8 catalyst although SEM-EDX shows that copper chromite has larger surface

compositions of copper elements (given in Table 7-5). Considering that both copper

surface area and copper dispersion may promote the hydrogenation reactions, it is

expected that rCu8 can perform better for the methyl formate hydrogenation reaction

under our investigated temperatures and pressures.

Table 7-4. Physicochemical properties of copper chromite and Cu8 catalysts


(m2/g)

Pore volume (cm3/g)

Cu surface area b

(m2/g)

Cu dispersionb

(%) Before reduction

After reduction a

Before reduction

After reduction a


183

Copper chromite

41.53 136.5 0.16 0.21 51.62 59.9

Cu8 147.1 144.0 0.22 0.25 47.25 80 a After reduction at 533 K. b Calculated from N2O dissociative adsorption.

Table 7-5. Surface composition of the catalyst

Sample Cu Zr Mg Al Cr O rCopper chromite

48.0% - - - 28.4% 23.6%

rCu8 21.80% 3.24% 14.08% 7.28% - 54.37%

7.4.2 CATALYTIC PERFORMANCE

HYDROGENATION AT 384 K

Two catalysts were used for the hydrogenation reaction of methyl formate at 384 K. The

experimental results are shown in Figure 7-6. It is found that the methanol production

rate and hydrogen consumption rate are nearly the same using these two catalysts.

However, the copper contents in each catalyst are significantly different as shown in

Table 7-5, but the copper surface area of two catalysts are nearly the same, where the

catalytic performance due to a slightly low copper surface area of rCu8 may be

compensated by its high copper dispersion. Based on the active copper component in the

catalysts, the STY was determined and tabulated in the Table 7-6. The STY of the rCu8

catalyst is double than that of copper chromite. This indicates that the reaction rate may

not depend on solely the copper contents but also depends on the copper surface area

and the copper dispersion in the catalysts. In addition, according to the literature [9],

[166], the addition of ZrO2 may also increase the adsorption of atomic hydrogen via the

‘spillover’ effect, thus increasing the reaction rate.


184

Figure 7-6. Amount of methanol and H2 in the reactor with two catalysts system. Operating conditions: Ptotal = 3.2 MPa, T = 384 K, stirrer speed: 800 rpm, catalyst loading: 16 g/L

Table 7-6. Space time yield of the catalysts at 384 K

Catalysts 𝑚𝑒𝑡ℎ𝑎𝑛𝑜𝑙 𝑔

𝑎𝑐𝑡𝑖𝑣𝑒 𝑐𝑎𝑡𝑎𝑙𝑦𝑠𝑡 𝑔 𝑖𝑛 𝑡ℎ𝑒 𝑓𝑖𝑟𝑠𝑡 90 𝑚𝑖𝑛𝑠

Copper chromite 9.03 Cu8 19.70

HYDROGENATION AT 370 K

These two catalysts were further tested at a lower temperature of 370 K with the same

initial total pressure to check if both catalysts perform comparably at other moderate

temperatures. The pressure profiles as well as the results can be found in Figure 7-7 and

Figure 7-8. It is interestingly observed that at the lower temperature, the developed rCu8

catalytically performs much better than the copper chromite. The modified STY listed in

Table 7-7 show that better performance is provided by the Cu/ZrO2-HTC catalysts

compared to that of copper chromite. Using the developed catalyst system at 370 K, the

hydrogenation reaction takes only 13 hours to reach equilibrium, compared to the 40

hours of using the commercial copper chromite. The slope of pressure decline using the

new catalyst system is larger than that of using the commercial cooper chromite at the

beginning, indicating that the immediate rate of the hydrogenation reaction is also

accelerated by using the developed rCu8 catalysts. It is speculated that at lower

temperatures, the copper dispersion plays a more important role to accelerate the


185

hydrogenation reaction. The existence of ZrO2 may also facilitate atomic hydrogen to be

adsorbed onto the surface of catalysts at low temperatures.

Figure 7-7. The total pressure profiles from two catalysts. Operating conditions: Ptotal = 3.2 MPa, T = 370 K, stirrer speed: 800 rpm, catalyst loading: 16 g/L

Figure 7-8. Amount of methanol and H2 in the reactor with two catalysts system. Operating conditions: Ptotal = 3.2 MPa, T = 370 K, stirrer speed: 800 rpm, catalyst loading: 16 g/L

Table 7-7. Space time yield of the catalysts at 370 K

Catalysts 𝑚𝑒𝑡ℎ𝑎𝑛𝑜𝑙 𝑔

𝑎𝑐𝑡𝑖𝑣𝑒 𝑐𝑎𝑡𝑎𝑙𝑦𝑠𝑡 𝑔 𝑖𝑛 𝑡ℎ𝑒 𝑓𝑖𝑟𝑠𝑡 3 ℎ𝑜𝑢𝑟𝑠

Copper chromite 6.94 Cu8 25.46


186

7.5 CONCLUSIONS

The novel Cu/ZnO/ZrO2-HTC catalysts were prepared using a simple co-precipitation

method and tested for the hydrogenation reaction of methyl formate to produce

methanol. The roles of each element in the catalytical system were identified by

comparing different compositions of the catalytic system. The catalytic performance of

different ratios of Cu/ZnO in the Cu/ZnO/ZrO2-HTC catalytic system was evaluated at 384

K and a pseudo-STY (STYPS) term was used to compare the reaction rate. rCu8 was

selected as the best catalytic system, and was further compared with the copper chromite,

which is a commercialised catalyst used for the hydrogenation reaction. Based on the

study, the following conclusions can be drawn:

(1) Copper species Cu0 and Cu+ are the active catalytic sites in the catalyst to promote

the hydrogenation reaction.

(2) The reaction rates can be significantly improved by incorporating the

Cu/ZnO/ZrO2 onto the HTC. HTC cannot only increase the total surface area of

catalysts, but it also helps increase copper dispersion on the surface of catalysts.

(3) The addition of ZnO and ZrO2 can increase the reaction rate to some extent,

depending on the catalyst compositions.

(4) 99.9 % selectivity can be obtained for all the prepared catalysts and a maximum

STYPS of 4.3 g methanol/g catalysts can be achieved when rCu8 is used.

(5) The STYPS is linearly correlated with the specific copper surface area in the

hydrogenation reaction, which has direct effect on the hydrogenation reaction

rate.

(6) The copper contents in the catalyst may not be a determining factor to evaluate

the catalytic performance of the hydrogenation reactions. The copper surface area


187

and copper dispersion resulting from ZrO2 play significant roles in promoting the

hydrogenation reactions. The possible reasons might be it has profound hydrogen

spillover effects which facilitate the transport of atomic hydrogen on the catalyst

surface, thus increasing the reaction rate.

(7) Through characteristic analysis and experimental hydrogenation reactions for all

prepared catalytic systems, rCu8 is the best catalyst. The catalysts were further

compared with a conventional hydrogenation catalyst, copper chromite. The rCu8

shows a better catalytic performance at moderate temperatures. The reaction rate

catalyzed by rCu8 is three times faster than that using copper chromite at 370 K.

A possible reason is that it enhances the hydrogen adsorption step on the catalyst

surface due to the existence of ZrO2.

Conclusions and recommendations

188

CHAPTER 8 CONCLUSIONS AND RECOMMENDATIONS

In the present study, a thorough literature review on possible effective methanol

synthesis pathways was conducted and presented in chapter 2, and a methanol synthesis

via methyl formate at moderate temperatures and pressures was selected for further

study. Literature review on solubility of reactant gases in reactant liquids were also

included in chapter 2. In chapter 3, materials and methodologies used and applied in the

study were explained. In chapter 4, an experimental apparatus was designed and set up

to study the vapour-liquid equilibrium of four systems containing CO-methanol, CO-

methyl formate, H2-methanol and H2-methyl formate. The results show that the

solubilities of CO and H2 in methanol and methyl formate increase with increasing

temperature and pressures possibly due to the endothermic processes and strong bond

interaction between gas molecules. A phi-phi approach using Peng-Robinson equations

of states was used to regress the experiments, and the binary interaction parameter 𝑘𝑖𝑗

for the four systems were regressed from the experimental data. The results indicate that

the binary interaction parameters of CO in methanol and methyl formate are temperature

independent, whereas the binary interaction parameters of H2 in methanol and methyl

formate are the function of temperature. In chapter 5, both carbonylation reaction and

hydrogenation reaction were preliminary studied and it was found that the

hydrogenation reaction is the rate determining step; therefore, the effects of agitation

speed, catalyst loadings and temperatures on the hydrogenation reaction rates using

copper chromite, a commercial catalyst, were investigated. A possible reaction

mechanism was proposed to evaluate the reaction kinetics parameters using the least

squares minimisation in MATLAB. Results show that the hydrogen adsorption and

dissociation is the slowest elementary step with the smallest reaction rate constant. This

Conclusions and recommendations

189

mechanism was further validated using the evaluated parameters to predict the

experiments conducted at three different pressures. In chapter 6, a novel catalyst system

consisting of copper, zinc oxide, zirconium oxide and hydrotalcite-like compounds was

designed and characterised using a number of techniques, including TPR, XRD, SEM, TGA,

BET, TPD-CO2, XPS and AES. The catalytic performance of the catalysts on the

hydrogeneration reaction was evaluated and presented in chapter 7. Results show that

the Cu/ZrO2-HTC provides the best performance. A comparison between the Cu/ZrO2-

HTC and copper chromite was thereafter conducted. It is found that the Cu/ZrO2-HTC can

improve the hydrogenation reaction rate three times at moderate temperatures, which

may because the addition of ZrO2 facilitates the transport and adsorption of atomic

hydrogen on the catalyst via the ‘spillover’ effect.

Based on the current discoveries, further experiments on the modifications of the novel

catalysts should be conducted, such as varying the ratio of copper and zirconium,

changing the hydrotalcite-like compounds percentage in the catalyst. In addition, other

catalysts preparation methods, such as sol-gel method, can be employed to compare the

catalysts performance on the hydrogenation reaction. Moreover, the investigation of

integrating the carbonylation reaction with the hydrogenation reaction using this novel

catalyst is worthwhile to check if this novel catalyst can be compatible with the different

catalyst used for the carbonylation reaction, and the total improvement can be made for

the two-step methanol synthesis process using the novel catalyst.

In future, such novel catalysts composed of copper, zinc and zirconium with optimised

ratio can be potentially used in the methanol synthesis industry at moderate operating

conditions to reduce the energy consumption significantly and further save operating

costs.

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Appendices

214

CHAPTER 10 APPENDICES

Appendix A Mass Spectrometry (MS) calibration of hydrogen using 5.4% H2/Ar

Table A. Calibration result of hydrogen gas via consecutive 20 repeat pulse injections.

Peak Area (intensity*s)

Temp. of pulse injection loop (K)

Pressure of pulse injection loop (kPa)

Dosing volume (cm3[STP])

moles (mmol)

Correction factor (CF) (mmol/intensity·s)

2.00208 × 10-10 301.9 100.67 0.0489 0.001960289 9.79126129 × 106

1.95767 × 10-10 301.9 100.67 0.0489 0.001960289 1.00133773 × 107

1.93217 × 10-10 301.9 100.66 0.0489 0.001960094 1.01445220 × 107

1.95508 × 10-10 301.9 100.66 0.0489 0.001960094 1.00256466 × 107

1.94655 × 10-10 301.9 100.65 0.0489 0.001959899 1.00685798 × 107

1.96778 × 10-10 301.9 100.65 0.0489 0.001959899 9.95995178 × 106

1.98453 × 10-10 301.9 100.62 0.0489 0.001959315 9.87294331 × 106

1.98251 × 10-10 302.0 100.62 0.0489 0.001958667 9.87973206 × 106

1.94853 × 10-10 302.0 100.59 0.0489 0.001958083 1.00516565E+07

2.01624 × 10-10 302.1 100.61 0.0488 0.00195382 9.69041586 × 106

1.99345 × 10-10 302.0 100.58 0.0488 0.001953884 9.80152129 × 106

1.98714 × 10-10 302.0 100.58 0.0488 0.001953884 9.83264521 × 106

2.01624 × 10-10 302.0 100.57 0.0488 0.00195369 9.73253693 × 106

1.99345 × 10-10 302.1 100.57 0.0488 0.001953044 9.80300868 × 106

2.00401 × 10-10 302.1 100.56 0.0488 0.001952849 9.74470896 × 106

1.97917 × 10-10 302.1 100.55 0.0488 0.001952655 9.86603082 × 106

2.01402 × 10-10 302.1 100.55 0.0488 0.001952655 9.69531197 × 106

2.02640 × 10-10 302.1 100.53 0.0488 0.001952267 9.63416318 × 106

1.99257 × 10-10 302.1 100.53 0.0488 0.001952267 9.79773271 × 106 9.74830967 × 106

Appendices

215

Appendix B. Mass Spectrometry (MS) calibration of CO2 using 4.99% CO2/He

Table B. Calibration result of carbon dioxide gas via consecutive 20 repeat pulse injections.

Peak Area (intensity·s)

Temp. of pulse injection loop(K) Pressure of pulse injection loop(kPa)

Dosing volume (cm3[STP])

moles (mmol)

Correction factor (CF) (intensity·s/mol)

7.33415 × 103 302.6 101.06 0.0453 0.001818793 4.032427 × 106

7.35292 × 103 302.5 101.05 0.0453 0.001819214 4.041811 × 106

7.31426 × 103 302.6 101.05 0.0453 0.001818613 4.021889 × 106

7.33317 × 103 302.6 101.06 0.0453 0.001818793 4.031888 × 106

7.33606 × 103 302.6 101.05 0.0453 0.001818613 4.033876 × 106

7.28585 × 103 302.7 101.05 0.0452 0.001813999 4.016457 × 106

7.38955 × 103 302.7 101.06 0.0452 0.001814179 4.073220 × 106

7.48585 × 103 302.8 101.07 0.0452 0.00181376 4.127256 × 106

7.31786 × 103 302.8 101.05 0.0452 0.001813401 4.035435 × 106

7.40583 × 103 302.8 101.07 0.0452 0.00181376 4.083138 × 106

7.33596 × 103 302.8 101.06 0.0452 0.001812982 4.046351 × 106

7.41364 × 103 302.9 101.06 0.0452 0.001812982 4.089197 × 106

7.33775 × 103 302.9 101.06 0.0452 0.001812982 4.047338 × 106

7.27024 × 103 303.0 101.08 0.0452 0.001812742 4.010631 × 106

7.24310 × 103 303.0 101.06 0.0452 0.001812384 3.996450 × 106

7.18381 × 103 303.1 101.07 0.0452 0.001811965 3.964651 × 106

7.25832 × 103 303.0 101.08 0.0452 0.001812742 4.004055 × 106

7.32010 × 103 303.0 101.09 0.0452 0.001812922 4.037737 × 106

7.23965 × 103 303.1 101.09 0.0452 0.001812324 3.994678 × 106

7.20385 ×103 303.0 101.09 0.0452 0.001812922 3.973614 × 106 4.002009 × 106

Appendices

216

Appendix C. Multi-level calibration of liquid samples in GC

Calibration of methanol and methyl formate

Table C. Mutil-level Calibration result of m ethanol and methyl formate mixture.

Level of calibration Methyl formate Methanol Retention Time: 1.742 min Retention time: 2.298 min Amount (µg/µl)

Area Amount (µg/µl)

Area

1 1.940 486.16 1.579 601.29 2 3.881 995.52 3.159 1196.4 3 4.851 1192.4 3.948 1525.2 4 5.821 1401.1 4.738 1812.4 5 7.762 1836.6 6.317 2399.1 6 8.732 2052.6 7.107 2675.1 7 9.702 2238.6 7.896 3024.0

Figure C-1. Calibration curve of liquid phase methyl formate using FID in GC.

Figure C-2. Calibration curve of liquid phase methanol using FID in GC.

Appendices

217

Appendix D. The grade of the certified standard-spec gas from ScottTM for retention time

determination

Appendices

218

Appendix E. Single-pint calibration of gas samples in GC

Hydrogen Gas

Calibration results Entry Area (a.u.) Amount (µL) 1 31803.1 250 2 32047.5 250 3 31716.4 250 4 31852.1 250 Average 31856.0 250

Properties of H2 Properties Values Units Density 0.083 kg/m3 Purity 99.999 % Volume 250 µL Mass 20.75 µg

Nitrogen Gas

Calibration results Entry Area (a.u.) Amount (µL) 1 4069 250 2 4049 250 3 4035 250 Average 4051 250

Properties of N2 Properties Values Units Density 1.251 kg/m3 Purity 99.999 % Volume 250 µL Mass 312.75 µg

Appendices

219

Methane Gas

Calibration results Entry Area (a.u.) Amount (µL) 1 16243 250 2 16336 250 3 16245 250 Average 16274.67 250

Properties of CH4 Properties Values Units Density 0.6752 kg/m3 Purity 99.999 % Volume 250 µL Mass 168.8 µg

Carbon Dioxide Gas

Calibration results Entry Area (a.u.) Amount (µL) 1 1525.86 250 2 1511.32 250 3 1549.7 250 Average 1528.96 250

Properties of CO2 Properties Values Units Density 1.809 kg/m3 Purity 99.999 % Volume 250 µL Mass 452.25 µg

Appendices

220

Carbon Monoxide Gas

Calibration results Entry Area (a.u.) Amount (µL) 1 3432 250 2 3416 250 3 3408 250 Average 3418.67 250

Properties of CO Properties Values Units Density 1.251 kg/m3 Purity 99.995 % Volume 250 µL Mass 312.75 µg

Dimethyl Ether Gas


Properties of DME Properties Values Units Density 2.11 kg/m3 Purity 4.96 % Volume 12.4 µL Mass 26.16 µg

Appendices

221

Oxygen Gas


Properties of DME Properties Values Units Density 2.11 kg/m3 Purity 21 % Volume 12.4 µL Mass 26.16 µg

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Author/s:

WU, Fan

Title:

Investigation of efficient two-step methanol synthesis processes at moderate pressures and

temperatures

Date:

2018

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