Catalysis in Flow: Monoalkylation Of

CATALYSIS IN FLOW: MONOALKYLATION OF

AMMONIA WITH ALCOHOLS

SUBMITTED IN PART FULFILMENT OF THE

REQUIREMENTS FOR THE DEGREE OF DOCTOR OF

PHILOSOPHY

ANDREW YUK KEUNG LEUNG

MARCH 2019

DEPARTMENT OF CHEMICAL ENGINEERING IMPERIAL COLLEGE LONDON

To my dear mum, Elizabeth Leung Fu Wai Ling,

and dad, Dr Leung Lun

Declaration

This thesis is submitted to Imperial College London for the degree of Doctor of Philosophy. It is a record of research carried out between March 2014 to February 2019 by the author, under the supervision of Professor Klaus Hellgardt and Professor Mimi Hii. It is believed to be wholly original, except where the due acknowledgement is made and has not been submitted for any previous degree at this or any other universities.

Andrew Yuk Keung Leung

3rd March 2019

The copyright of this thesis rests with the author and is made available under a Creative Commons Attribution Non-

Commercial No Derivatives licence. Researchers are free to copy, distribute or transmit the thesis on the condition that they attribute it, that they do not use it for commercial purposes and that they do not alter, transform or build upon it. For any reuse or redistribution, researchers must make clear to others the licence terms of this work.

Abstract

This PhD thesis describes the Ni-catalysed alkylation of ammonia with alcohols to achieve high selectivity to primary amines. This work was broadly divided into two parts: (i) developing and understanding reaction conditions through batch reactions; and (ii) deploying the system into a flow reactor and understanding how the system works and fails.

Extensive screening of heterogeneous catalysts in the alkylation of ammonia with alcohols was carried out in batch reactors. It was found that a commercial

65 wt% Ni/Al2O3/SiO2 catalyst yielded the highest selectivity towards the primary amine. High selectivities of 99% with 7 different alcohols to their corresponding amines were achieved with moderate alcohol conversions, under the reaction conditions of 160 °C, 72 hours, 30 mL of 0.1 M alcohol in o-xylene, alcohol/Ni = 10 and anhydrous ammonia/alcohol = 7. An extruded version of this catalyst was created for the application in flow.

A flow reactor was then designed and constructed. Competitive formation of the nitrile side-product was suppressed when the catalyst was pre-reduced. Higher alcohol conversions and selectivities to the primary amines were achieved at 51

– 100% and 99% respectively. This improvement in performances is attributed to the minimisation of water accumulation. Over an extended continuous run for

78 hours, it was found that the carbon deposition on the catalyst resulted in the deactivation of the catalyst in flow.

Acknowledgements

I am very grateful to my supervisors Professor Mimi Hii (Department of

Chemistry) and Professor Klaus Hellgardt (Department of Chemical

Engineering) for their constant support and excellent supervision over the past four years. I could not have done this degree without their inspiration and encouragement!

I would also like to thank all the members of REaCT group as well as the Barton lab, past and present. With special thanks to John, Sergio, Fessehaye, Bhavish,

Suna, Lisa, Ben(s), Kane, Ilia, Isaac and Faye, for all the fun times we’ve had in and out of the labs. Thank you to Susi, Severine, Graham, Anthony for making my time in the department run smoothly. A special thank you to Patricia for her patience.

Lastly, thank you to my family for their support, especially my parents and brothers for their understanding and patience for what I do.

Nomenclature

Symbol or Meaning Usual units

Abbreviations

1H-NMR Proton Nuclear Magnetic Resonance

1° Primary (amine)

2° Secondary (amine)

3° Tertiary (amine)

A Frequency factor (in Arrhenius Equation)

atm Standard atmosphere (1.01325 bar)

BET Brunauer-Emmett-Teller

b.p. Boiling point K

d Diameter m

DFT Density Functional Theory

-1 ΔHf Enthalpy of formation kJmol

-1 ΔHr Enthalpy of reaction kJmol

-1 Ea Activation energy kJmol

EDX Energy Dispersive X-ray

ENRTL-RK Electrolyte Non-Random Two-Liquid with Redlich-

Kwong equation of state model

EOS Equation of State

eq. Equation

equiv. equivalent(s)

ΔG Change of Gibbs free energy kJmol-1

GC Gas Chromatography

GC-FID Gas Chromatography – Flame Ionisation Detector

GC-MS Gas Chromatography – Mass Spectrometry

GS General Solver (for Berkeley Madonna, a modelling

software)

i- iso-

ICP Inductively Coupled Plasma elemental analysis

I.D. Inner Diameter m

K Equilibrium constant No units

k Reaction rate constant s-1

M Molar mol dm-3

O.D. Outer Diameter m

o- ortho-

p- para-

R Universal Gas Constant J K-1 mol-1

RK4 Runge-Kutta method (fourth order)

RTP Room Temperature Pressure

rpm Revolutions per minute min-1

SR-Polar Schwartzentruber-Renon-Polar

ΔS Entropy kJ K-1

TEA Thermodynamics for Engineering Applications

TEM Transmission Electron Microscopy

TON Turnover Number mol/mol

TPD Temperature Programmed Desorption

VSEPR Valence Shell Electron Pair Repulsion Theory

All other abbreviations are described in the texts as they appear.

List of Schemes, Figures and Tables

Schemes

Scheme 1.1. Examples of amines with different applications.

Scheme 1.2. Examples of primary amines used in the pharmaceutical and polymer industries

Scheme 1.3. Model reaction to test the catalyst efficiency in the N-alkylation of ammonia with benzyl alcohol.

Scheme 2.1. (From left to right) Ammonia, and primary, secondary and tertiary amines, where R = alkyl or aromatic groups

Scheme 2.2. Reaction scheme of the Haber-Bosch process with typical industrial conditions

Scheme 2.3. Molecular structure of ammonia. The lone electron pair repels the shared electron pairs, causing a slight decrease in bond angles.

Scheme 2.4. Ethene hydrogenation with a nickel catalyst at 150 ⁰C.

Scheme 2.5. Catalytic conversions of poisonous gas molecules to less harmful ones.

Scheme 26. Ziegler-Nata polymerisation using a homogeneous Ti catalyst.

Scheme 2.7. (Top) Ziegler-Nata polymerisation termination with the β-elimination from the polymer chain. (Bottom) Ziegler-Nata polymerisation termination with the β- hydrogen elimination reaction.

Scheme 2.8. General scheme of the Wacker process.

Scheme 2.9. Modern formulation of the catalytic cycle of the Wacker process.

Scheme 2.10. Summary of primary alkyl amine synthesis methods, where common examples of reactants are included.

Scheme 2.11. General reaction scheme for alkyl alcohol to a halide, and subsequently substitution by an amine to form an alkyl amine.

Scheme 2.12. Gabriel synthesis. An alcohol was converted to the halide and reacted with phthalimide to yield a primary amine.

Scheme 2.13. 3,4-diphenylmaleic anhydride as a catalytic Gabriel reagent in an amine substitution reaction.

Scheme 2.14. Proposed mechanistic pathway of the Cu-catalysed insertion of carbene into the N-H bond of phthalimide for the amine substitution reaction.

Scheme 2.15. Example of aryl primary amine production by Buchwald-Hartwig cross- coupling using a homogeneous Fe catalyst [72].

Scheme 2.16. Metal-free amination of alkyl boronic acids [73].

Scheme 2.17. A general scheme of hydroamination reaction between ammonia and an alkene.

Scheme 2.18. An early example of hydroamination using sodium as a reactant [75].

Scheme 2.19. Industrialised hydroamination of isobutylene to t-butylamine [78].

Scheme 2.20. Homogeneous catalytic hydroamination of allyl groups with ammonia using a homogeneous Au catalyst [83].

Scheme 2.21. Formal one-pot, two-step hydroamination of olefins using a homogeneous Pd/Ir dual metal tandem catalyst system [84].

Scheme 2.22. Enantioselective hydroamination of olefins by biocatalysts to a carboxylic acid [86].

Scheme 2.23. A general scheme of hydromethylation of olefins to produce primary amines. Note the increase in the length of the compound by one carbon [89].

Scheme 2.24. Top: reaction scheme for the hydroaminomethylation of pentene, butene and propene. Bottom: Ligands used by Zimmermann et al. with Rh and Ir to achieve selective hydroaminomethylation to primary amines [90].

Scheme 2.25. Hydroaminomethylation of limonene by Behr et al. [92].

Scheme 2.26. A general scheme of nitrile reduction with common metal catalysts and

H2 as the reductant.

Scheme 2.27. Potential side reactions of nitrile reduction.

Scheme 2.28. Nitrile reduction with a homogeneous Fe complex by Lange et al. [99].

Scheme 2.29. Selective reduction of nitriles using a cobalt phosphine catalyst by Adam et al. [100].

Scheme 2.30. Flow scheme of selective hydrogenation of nitriles to primary amines catalysed by a supported Pd catalyst, reported by Saito et al. [102].

Scheme 2.31. Boron-catalysed silylative reduction of nitriles to primary amines [103].

Scheme 2.32. Example of reductive amination of an aldehyde to form a primary amine.

Scheme 2.33. Potential pathways for reductive hydrogenation.

Scheme 2.34. Scheme of reductive amination using CO or H2 as reductants.

Scheme 2.35. Amine dehydrogenases catalyse the reductive amination of ketones and aldehydes to chiral amines. Conditions: AmDH (30 – 130 µM), ammonium formate buffer (1.005 M, pH 8.5), T = 30 °C, agitation on orbital shaker (190 rpm), 24 – 48 hours.

Scheme 2.36. Reductive amination with Ni-Al alloy under ultrasound.

Scheme 2.37. Alcohol preparations before conversion to primary amine.

Scheme 2.38. Mechanism of the hydrogen borrowing technique [113].

Scheme 2.39. One of the first reported examples of alkylation of amines with alcohols by Winans and Adkins [115].

Scheme 2.40. Early homogeneous reactions using a homogeneous Ru catalyst in the hydrogen borrowing cycle [117].

Scheme 2.41. Structure of acridine-based pincer complex [RuHCl(A-iPr-PNP)(CO)] used in the alkylation of ammonia via the hydrogen borrowing cycle [118].

Scheme 2.42. Structure of pincer ligand used by Pingen et al. in the homogeneous Ru catalytic system in the alkylation of ammonia via the hydrogen borrowing cycle. [119]

Scheme 2.43. Two-enzyme cascade regime for the hydrogen borrowing alkylation of

+ amines. Conditions. 30 C, 24 – 48 hours, 5 mol% enzymes, 2M NH4 /NH3 [130].

Scheme 2.44. Nucleophilicities of ammonia, a primary amine and a secondary amine in increasing order.

Scheme 3.1. Mechanism of hydrogen borrowing technique

Scheme 4.1. Model of the N-alkylation of NH3 with BnOH with labelled kinetic constants

Scheme 5.1. Scheme for hydrogen borrowing cycle of NH3 alkylation with BnOH, where R = Ph.

Scheme 5.2. Primary amine formation from BnOH+NH3.

Scheme 5.3. Expected Bn2NH and Bn3N formation from BnOH+NH3 reaction.

Scheme 5.4. Possible reaction products with BnOH and NH3 are BnNH2, Bn2NH and

Bn3N.

Scheme 5.5. Reaction A: overall reaction scheme for BnOH+NH3.

Scheme 5.6. Reaction 1: The first step of reaction A, where BnOH reacts with NH3.

Scheme 5.7. Reaction 2: The second step of reaction A, where BnOH reacts with

BnNH2.

Scheme 5.8. Reaction 3: The third step of reaction A, where BnOH reacts with

Bn2NH.

Scheme 5.9. Hydrogen borrowing cycle with emphasis on the condensation step.

Scheme 5.10. Examples of N-alkylation of NH3 with BnOH that resulted in different selectivities.

Scheme 5.11. Proposed scheme of formation of aldehyde due to an incomplete hydrogen borrowing cycle.

Scheme 5.12. Reaction scheme of BnOH and aniline forming secondary amine.

Scheme 5.13. Disproportionation reaction of benzyl alcohol to toluene and benzaldehyde.

Scheme 5.14. A three-reaction kinetic model in the NH3 N-alkylation reaction with BnOH.

Scheme 5.15. Hydrogen borrowing cycle of the reaction between BnNH2 and BnOH.

Scheme 5.16. A four-reaction kinetic model to describe the reaction between the N- alkylation of NH3 with BnOH (reaction A).

Scheme 5.17. A recap of the three-reaction kinetic model used for reaction A.

Scheme 5.18. A recap of the four-reaction kinetic model used for reaction A.

Scheme 5.19. Hydrogen borrowing cycle with emphasis on the condensation step.

Scheme 5.20. The two-reaction model used in calculating the kinetic constants for the temperature studies of reaction A.

Scheme 5.21. Menshutkin alkylation of ammonia, where R is the leaving group and X is the halide.

Scheme 5.22. Steric effects of a 2-methoxyl benzaldehyde (left) compared to a 3- methoxyl benzaldehyde during the condensation reaction with NH3.

Scheme 5.23. Resonance effect of a 3-picolyl benzaldehyde (left) compared to the absence of the effect in a 2-picolyl benzaldehyde.

Scheme 5.24. Incomplete hydrogen borrowing cycle of 2-octanol, resulting in ketones but no imine or target amines.

Scheme 5.25. Formal one-pot, two-step hydroamination of olefins using a homogeneous Pd/Ir dual metal tandem catalyst system [84].

Scheme 5.26. Hydroaminomethylation of limonene by Behr et al. [92].

Scheme 6.1. A simplified scheme of a generic plug flow reactor.

Scheme 6.2. Possible pathway of benzonitrile formation from the hydrogen borrowing cycle, when the catalyst was not fully reduced.

Scheme 6.3. Steric effects of a 2-methoxybenzaldehyde (left) compared to a 3-methoxybenzaldehyde during the condensation reaction with NH3.

Scheme 6.4. Flow scheme of selective hydrogenation of nitriles to primary amines catalysed by a supported Pd catalyst, reported by Saito et al. [102].

Scheme S1. Synthetic route of Lisinopril [3].

Scheme S2. Three-reaction model of the N-alkylation of NH3 with BnOH with labelled kinetic constants.

Scheme S3. Four-reaction model of the N-alkylation of NH3 with BnOH with labelled kinetic constants.

Figures

Figure 4.1. Batch reactor combined with the condensation tube.

Figure 4.2. Condensation tube to introduce anhydrous NH3 into the batch reactor.

Figure 4.3. Schematic of the customised flow reactor

Figure 4.4. Labelled photograph of the flow reactor

Figure 4.5. Schematic of the flow reactor configuration for catalyst reduction

Figure 4.6. Calibration curves for Gilson 305 HPLC pumps, 25 cc (left) and 5 cc pump heads

Figure 4.7. Dispersion profile of the flow reactor at 0.5 mL/min

Figure 4.8. XRD patterns for Ni catalysts used in batch reactions

Figure 4.9. The process of making the 36 wt% Ni Catalyst extrudites (Catalyst 2)

Figure 4.10. Simulation flow sheet in COFE with NH3 and o-xylene as the components.

Figure 4.11. Physical properties available from the COFE simulation. The “heated” flow stream indicated the mole fractions.

Figure 4.12. COFE simulated phase diagram of different stoichiometries of NH3/BnOH solutions. Conditions 160 C, pressure = 0 – 60 bar. From left to right: NH3/BnOH = 1, 2, 3, 4, 5, 7, 10.

Figure. 4.13. Aspen simulation using the SR-Polar method: Phase diagram of NH3/o- xylene mixture at 60 bar. The black curve represents the dew point curve and blue the bubble point curve. These curves separate the liquid, vapour and vapour + liquid phases.

Figure 4.14. In-built chemical equations function interface, where the reaction steps and initial concentrations are defined.

Figure 4.15. Example of a Berkeley Madonna generated fit. Using data from a batch reaction with a stoichiometry of BnOH/NH3 = 1.

Figure 4.16. Temperature programmed desorption equipment setup flow diagram. The components shown in the images are listed in Table 4.5.

Figure 4.17. Calibration graph for Bn2NH. Concentrations of Bn2NH solutions were analysed by GC-FID (x-axis) against (area of the component)/(area of the standard) (y- axis)

Figure 5.1. XRD patterns for the commercial catalysts, 65 wt% Ni/Al2O3/SiO2 and 16 wt% NiMo/Al2O3.

Figure 5.2. Reaction profiles of BnOH and NH3 catalysed by 65 wt% Ni/Al2O3/SiO3 for 48 hours (Top) and 1 wt% Au/TiO2 for 4 hours (Bottom). Kinetic profiles were generated by Berkeley Madonna, where the experimental data are shown as data points and the calculated results are represented by solid curves.

Figure 5.3. The molar ratio of H2O/NH3 was plotted against conversion.

Figure 5.4. Arrhenius plot using data from anhydrous NH3 alkylation with benzyl alcohol at 80 – 160 °C.

Figure 6.1. Schematic of the NH3 condensation cylinder.

Figure 6.2. Gas chromatograph of a sample from a flow reaction containing PhCN

Figure 6.3. Data collected from flow Reaction A. Benzonitrile (PhCN, black squares) is observed as the major product initially. Conditions: initial BnOH = 0.6 M, 160 °C,

0.2 mL/min, NH3/BnOH = 7, 2.0 g Catalyst 2.

Figure 6.4. Reaction profile of flow reactions using reduced Catalyst 2. Conditions were

0.06 mL/min, 0.1 M BnOH in o-xylene, 160 °C, 60 bar, NH3/BnOH = 7, 2.0 Catalyst 2. The graph is showing results from the first 4 hours of the reaction.

Figure 6.5. Reaction profile of flow Reaction A using reduced Catalyst 2. Conditions were 0.06 mL/min, 0.1 M BnOH in o-xylene, 160 °C, 60 bar, NH3/BnOH = 7. The graph shows results for up to 14 hours.

Figure 6.6. Plot of conversion of 0.6 M BnOH and selectivity to BnNH2 under different flow rates. Conditions: 0.1 – 1 mL/min, 2.0 g Catalyst 2, 160 °C, 60 bar, 0.60 M BnOH in o-xylene, NH3/BnOH = 7. Conv. = Conversion; Sel. = Selectivity.

Figure 6.7. Plot of conversion of 0.2 M BnOH and selectivity to BnNH2 under different flow rates. Conditions: 0.1 – 1 mL/min, 2.0 g catalyst, residence time = 2.8 - 28 min,

160 °C, 60 bar 0.2 M BnOH in o-xylene, NH3/BnOH = 7, 2.0 g Catalyst 2.

Figure 6.8. Plot of conversion of 0.1 M BnOH and selectivity to BnNH2 under different flow rates. Conditions: 0.06 – 0.5 mL/min, 2.0 g catalyst, residence time = 2.8 – 47 min,

160 °C, 60 bar 0.1 M BnOH in o-xylene, NH3/BnOH = 7, 2.0 g Catalyst 2.

Figure 6.9. Plot of conversion of 0.1 M, 0.2 M and 0.6 M BnOH to BnNH2 under different flow rates. Conditions: 0.06 – 1 mL/min, 2.0 g catalyst, residence time = 2.8 –

47 min, 160 °C, 60 bar 0.1 M BnOH in o-xylene, NH3/BnOH = 7, 2.0 g Catalyst 2.

Figure 6.10. The TON and selectivity to BnNH2 of an extended experiment over 80 hours. Conditions: 0.06 mL/min, 2.0 g Catalyst 2, residence time = 93 min, 160 °C, 60 bar 0.10 M ROH in o-xylene, NH3/BnOH = 7.

Figure S1. Phase diagram of ammonia, from pressures between 0 to 50 bar. The plots are vapour pressures of the gas at different temperatures Lange's Handbook of Chemistry, 10th ed. page 1451 and 1468.

Figure S2. NH3 solubility in toluene at 120 °C, Taken from Bhattacharyya et al. [167].

Figure S3. Reaction profile of the synthesised 85 wt% Ni/Al2O3/SiO2 catalyst. (38 – 90 µm).

Figure S4. Reaction profile of the synthesised and reduced 85 wt% Ni/Al2O3/SiO2 catalyst (38 – 90 µm).

Figure S5. Reaction profile of the synthesised 85 wt% Ni/Al2O3/SiO2 catalyst (90 – 250 µm).

Figure S6. Reaction profile of the synthesised and reduced 85 wt% Ni/Al2O3/SiO2 catalyst (90 – 250 µm).

Figure S7. Example of a General Solver fit using multiple sets of experimental data.

Figure S8. Four-reaction model fit for BnOH+NH3+H2O, Table 5.6 Entry 1.

Figure S9. Three-reaction model fit for BnNH2+H2O, Table 5.5 Entry 2.

Figure S10. Four-reaction model fit for BnNH2+H2O, Table 5.6 Entry 2.

Figure S11. Three-reaction model fit for Bn2NH+H2O, Table 5.5 Entry 3.

Figure S12. Four-reaction model fit for Bn2NH+H2O, Table 5.6 Entry 3.

Figure S13. Three-reaction model fit for Bn3N+H2O, Table 5.5 Entry 4.

Figure S14. Four-reaction model fit for Bn3N+H2O, Table 5.6. Entry 4.

Figure S15. Two-reaction model fit for BnOH:NH3:H2O = 1:1:5

Figure S16. Three-reaction model fit for BnOH:NH3:H2O = 1:1:5, Table 5.8. Entry 1.

Figure S17. Two-reaction model fit for BnOH:NH3:H2O = 1:1:10

Figure S18. Two-reaction model fit for BnOH:NH3:H2O = 1:1:10. Table 5.8. Entry 2.

Figure S19. Two-reaction model fit for BnOH:NH3:H2O = 1:2:10

Figure S20. Three-reaction model fit for BnOH:NH3:H2O = 1:2:10. Table 5.8. Entry 3.

Figure S21. Three-reaction model fit for BnOH:NH3:H2O = 1:2:20. Table 5.8. Entry 4.

Figure S22. Two-reaction model fit for BnOH:NH3:H2O = 6:1:0

Figure S23. Three-reaction model fit for BnOH:NH3 = 6:1:. Table 5.8. Entry 6.

Figure S24. Two-reaction model fit for BnOH:NH3 = 1:3.

Figure S25. Two-reaction model fit for BnOH:NH3 = 1:3 at 80 °C

Figure S26. Two-reaction model fit for BnOH:NH3 = 1:3 at 120 °C

Figure S27. Two-reaction model fit for BnOH:NH3 = 1:3 at 160 °C

Figure S28. Universal kinetic constants fit back into the starting Ni experiment.

Figure S29. TEM images of 65 wt% Ni-Al2O3/SiO2. The bar chart shows catalyst particle size distribution, data collected by 182 particles using four images

Figure S30. TEM image of 85 wt%Ni/Al2O3/SiO2

Figure S31. XRD spectra for comparison of catalysts at different stages. Mt represents montmorillonite, a form of Al2O3/SiO2 present in bentonite clay

Figure S32. XRD patterns of synthesised 85 wt% Ni/Al2O3/SiO3 catalysts. The graphs are identical. The bottom graph is labelled for clarity.

Tables

Table 4.1: List of components used in the customised flow reactor and their description.

Table 4.2. NH3 alkylation with BnOH by synthesised and commercial Ni catalysts.

Table 4.3. NH3 alkylation with BnOH by synthesised (reduced) and commercial Ni catalysts

Table 4.4. Conditions of each stream in the COFE simulation

Table 4.5. Calculated individual flow rates of the NH3/o-xylene mixture

Table 4.6. Components of the Temperature Programmed Desorption equipment

Table 4.7. Retention times of common species detected by GC-FID

Table 5.1. Thermodynamics data predicted by ASPEN-plus with ENTRL-RK method.

Table 5.2. Preliminary batch results with commercial catalysts

Table 5.3. Kinetic constants of commercial Ni- and Au- catalysed NH3 reaction A

Table 5.4. Adjusted kinetic constants for commercial Ni- and Au- catalysed reaction A.

Table 5.5. Ni catalysed reactions of product amines with water. Initial concentrations of each substrate are listed in shaded boxes for clarity.

Table 5.6. Ni-catalysed reactions of product amines with water (set 1A). Kinetic constant values were fitted by using the three-reaction model (Scheme 5.17)

Table 5.7. Ni-catalysed reactions of product amines with water (Set 1B). Kinetic constant values were fitted by using the four-reaction model (Scheme 5.18).

Table 5.8. Conversions and selectivities from reaction A catalysed by Catalyst 1, using different BnOH:NH3:H2O ratios

Table 5.9. Kinetic constants from reaction A by Catalyst 1 in batch, using different

BnOH:NH3:H2O ratios. The kinetic constant values were fitted by Berkeley Madonna using the three-reaction model.

Table 5.10. General solvers calculated from reaction sets 1A and 2.

Table 5.11. Temperature study of reaction A using Catalyst 1.

Table 5.12. Results of N-alkylation of NH3 with different alcohols using Catalyst 2.

Table 6.1: Flow reaction results from 120 to 200 °C.

Table 6.2: Flow reaction results at different catalyst grades.

Table 6.3: Brunaun-Emmett-Teller (BET) specific surface area measurements of Ni catalysts of different sizes.

Table 6.4. Flow reaction results at different NH3/BnOH ratios.

Table 6.5: Fixed bed flow reactions results with various alcohols.

Table 6.6. Ni catalysts at different stages and their BET & TPD results.

Table S1. A prediction of the thermodynamic outcome from ASPEN, presented in mol%.

Table S2. Sample code for calculating the general solver on Berkeley Madonna.

Table S3: Ni catalysts at different stages and their BET & TPD results.

Table of Contents

Declaration ...... 3

Abstract ...... 5

Acknowledgements ...... 6

Nomenclature ...... 7

List of Schemes, Figures and Tables ...... 9

Table of Contents ...... 21

Project Motivation ...... 24

Chapter 1: Introduction ...... 25

1.1. Importance of Primary Amines ...... 26

1.2. Flow vs Batch Processes ...... 28

1.3. Conclusion ...... 30

Chapter 2: Literature Review ...... 31

2.1. Ammonia, The Nitrogen Source...... 31 2.2. Heterogeneous and Homogeneous Catalysis ...... 35 2.2.1. Heterogeneous Catalysis ...... 35

2.2.2. Homogeneous Catalysis ...... 42

2.2.3. Summary and The Future of Catalysis ...... 46

2.3. Primary Amine Production Methods ...... 47 2.3.1. Alkyl Halide Substitution ...... 48

2.3.2. Olefin Hydroamination ...... 52

2.3.3. Hydroaminomethylation ...... 56

2.3.4. Nitrile or Nitro-group Hydrogenation ...... 59

2.3.5. Reductive Amination ...... 63

2.3.6. Amine Alkylation with Alcohol via Hydrogen Borrowing Cycle ...... 68

2.3.7. Critical Comparison of the Amine Production Methods ...... 78

2.4. Conclusion ...... 79 Chapter 3: Aims and Objectives ...... 81

3.1. Aims ...... 81 3.2. Objectives ...... 82 Chapter 4: Experimental Protocol and Methodology .. 84

4.1. Introduction ...... 84 4.2. Materials ...... 84 4.3. Experimental...... 86 4.3.1. Batch Reaction ...... 86

4.3.2. Flow Reaction ...... 90

4.3.3. Catalyst Preparation and Preliminary Performances ...... 99

4.4. Calculations ...... 105 4.5. Computational Methods ...... 106 4.5.1. Phase Diagram Simulations ...... 106

4.5.2. Curve Fittings for Kinetic Studies ...... 114

4.6. Characterisation Techniques ...... 117 4.7. Conclusion ...... 131 Chapter 5: Batch Reactions ...... 132

5.1. Introduction ...... 132 5.2. Thermodynamics ...... 136 5.3. Preliminary Batch Experiments with Commercial Catalysts ...... 144 5.4. Kinetic Studies...... 152 5.4.1. Kinetic Models with Ni and Au Catalysed Reactions ...... 153

5.5. Reactions with Intermediates and Refined Kinetic Model with Ni ..... 159 5.6. Effect of Water ...... 139 5.7. Temperature Studies ...... 150 5.8. Batch Reactions with Different Alcohols ...... 154 5.9. Comparisons with Different Primary Amine Production Methods ..... 157

5.10. Conclusion ...... 159 Chapter 6: Flow Reactions...... 161

6.1. Introduction ...... 161 6.2. System Characterisation ...... 163 6.2.1. Flow Reactor Considerations and Specifications ...... 164

6.3. Preliminary Flow Reactions ...... 166 6.3.1. Unexpected Intermediate: Benzonitrile ...... 167

6.4. Conditions Study ...... 172 6.4.1. Reaction temperature ...... 172 6.4.2. Catalyst size ...... 174 6.4.3. Ammonia/Alcohol Molar Ratio ...... 177 6.4.4. Flow Rates ...... 179 6.4.5. Initial Alcohol Concentrations ...... 181 6.4.6. Summary for Reaction Conditions ...... 184 6.5. Flow Reactions with Different Alcohols ...... 185 6.6. Comparisons with Different Primary Amine Production Methods ..... 187 6.7. Extended flow reaction over time ...... 189 6.8. Catalyst deactivation ...... 191 6.9. Conclusion ...... 192 Chapter 7: Summary and Recommendations for

Future Work ...... 193

7.1. Summary...... 193 7.2. Recommendations for Future Work ...... 197 Appendix ...... 201

References ...... 231

Project Motivation

PhD motivation

This project was motivated by the following aspects:

• To provide a more effective step to synthesis amines from alcohols, preferably

using the hydrogen borrowing cycle.

• To become an expert in research, specifically in process engineering and

catalysis.

• To enjoy the knowledge sharing culture that is academic research.

• To learn transferable skills, for example research and collaboration, which are

key skills in many industries.

• To fulfil personal curiosity.

• To challenge myself.

Chapter 1: Introduction

Amines belong to a very important class of organic chemicals that are a part of our daily lives. They are nitrogen-containing compounds which are involved in the production of dyes, pharmaceuticals and agrochemicals and carbon dioxide capturing agents for use in oil refinery systems [1]. They are used in the production of many drugs and herbicides such as glyphosate, atrazine and lisinopril, which are among the most used

pharmaceutical chemicals in the world (Scheme 1.4) [2-3].[2][3] Glyphosate was once the most widely used herbicide in the US agricultural sector with 82,000 – 84,000 tonnes applied annually, but its use has decreased dramatically over the years due to weed resistance, contamination of surface water due to run-off and its human toxicity [4].

Atrazine is a herbicide that prevents pre- and post-emergence broadleaf weeds. It is the most widely used herbicide after glyphosate in the United States, with 35,000 tonnes used annually [5]. Lisinopril is a drug to treat high blood pressure and was the second most prescribed drug in the United States in 2017 [6]. Example of amines used in different industries are shown in scheme 1.1.

Scheme 1.1. Examples of amines with different applications.

1.1. Importance of Primary Amines

There are lower (C1-6) and higher (Cn>6) primary amines with differences in functionality and utility. Smaller amines are generally used as building blocks, while higher amines

(C6 or above) are particularly valuable, as they can be used to produce additives in personal-use products, such as polymers (Nylon-6,6) [7], for agrochemicals and in neuroactive pharmaceutical ingredients, such as memantine for treatment of

Alzheimer’s disease [8], Pregabalin for treatment of anxiety disorder [9] and the precursor to SR58611A, an antidepressant [10] (Scheme 1.5).

Scheme 2.2. Examples of primary amines used in the pharmaceutical and polymer industries

The smaller amines can be made by synthesised gas phase catalysis [11], while the larger amines require different approaches such as reductive amination [12] or the

Gabriel synthesis [13]. Among the many methods of synthesising primary amines, one example that is becoming more popular is the hydrogen borrowing technique due to its atom-efficiency as well as water being the only by-product. A common model reaction to test the efficiency of a catalyst involves benzyl alcohol (BnOH) and amines as seen in many examples [14-17] (Scheme 1.6).

Scheme 1.3. Model reaction to test the catalyst efficiency in the N-alkylation of ammonia with benzyl alcohol.

Benzyl alcohol is commonly used as a model reaction for its availability and is a staple for alcohol to amine conversions. The use of BnOH allows easier comparison of this reaction with other literature, as many research groups had selected BnOH as the model alcohol in their work. However, the clear drawback is that BnOH is a simple aryl alcohol with no other functional groups, so reactivity with other functional groups cannot be tested.

A detailed literature review in Chapter 2 will explore the production methods of primary amines.

1.2. Flow vs Batch Processes

Flow chemistry has been flourishing over the past decade as its concept aligns well with the philosophy of green chemistry [18]. As stated in the roundtable at the green chemistry conference, the use of green chemistry is recommended [19].

The use of small continuous flow reactors to perform chemical synthesis in laboratories in particular has become more popular, due to the following advantages over traditional batch reactors [20]:

• improved heat and mass transfer due to the high surface area contact with the

catalyst,

• elimination of free headspace,

• no accumulation of reactive intermediates or hazardous materials,

• a steady-state operation which makes for better quality control, ease of

automation and stability of the overall process, as well as lower operating cost

(lower operational expense (OPEX)) [21], and

• continuous process which allows simple reaction scale-up, addition of

sequential synthetic reactions with independent reaction conditions and in-line

purification, analysis or monitoring methods.

Due to these advantages, the use of flow reactors for industrial applications as well as small scale productions is expected to grow in the years to come [22].

On the other hand, some advantages of batch reactors over flow reactors may include:

• generally lower set up cost (lower OPEX),

• ease in maintenance and inspections of equipment between reactions,

• faster development cycles for new products, and

• generally better mixing of reactants.

The pros and cons of each process must then be examined before deciding on the better suited method for this project. Take the following case as an example: a Novartis-MIT continuous pilot plant produces 100 g/hr of aliskiren using two synthetic steps and is able to produce Tablets containing 112 mg of free aliskiren. The continuous reactor used in this work as a reactor volume of 0.7 L, which can prepare 0.8 tons/year of the active pharmaceutical ingredient (API). In the commercial scale of 188 tons/year, a reactor volume of only 136 L is required, which is only one tenth the size of the actual batch reactor volume, which is 1500 L. Furthermore, the flow reaction operates in solventless conditions and is completed in 1 hr, compared to the batch reactions which require 48 hours in refluxing conditions. These numbers suggest that the automated flow process is a more environmentally friendly process with a smaller overall footprint [23].

In this project, ammonia was heavily used across all reactions. As ammonia is a toxic and highly volatile reagent [24], the use of flow reactors should be encouraged. This is because a flow system can ensure the safety of the work environment as only small amounts of reagents are heated at any instant and with great reaction control, minimising the users’ exposure to the toxic gas. The stoichiometry of ammonia can be easily adjusted by different flow rates, as well as the overall reaction times.

Our group had previously demonstrated the borrowing hydrogen approach in flow, where the N-alkylation of primary and secondary amines by alcohols using Au as a catalyst was performed using a commercial flow reactor [25]. Other scalable solutions to organic synthesis through the implementation of catalytic processes in flow were achieved, such as the observation of Pd catalyst leaching during Suzuki-Miyaura cross- coupling reactions using a custom-made plug flow reactor [26], and the effect of O2 on the catalytic activity of a Ru catalyst during the aerobic oxidation of an alcohol was observed online using a plug flow differential reactor [27]. This project aims to extend the approach to the synthesis of primary amines using ammonia as the feedstock.

1.3. Conclusion

It is important to establish new methods for the synthesis of primary amines, and alternative methods by using cheaper and more available metals as catalysts should be explored. Moreover, the production must be clean and sustainable with the least amount of waste possible. Techniques in flow chemistry can achieve this endeavour. However, challenges still exist in achieving high selectivity to primary amine, an aspect addressed by many studies.

The next chapter discusses recent developments in the field of primary amine production methods, including catalysed hydrogen borrowing cycle, with examples provided for heterogeneous and homogeneous catalysts, and the pros and cons of different production methods are also discussed in detail.

Chapter 2: Literature Review

This chapter first discusses ammonia as a nitrogen source for the chemical industry, followed by discussing the merits and drawbacks of heterogneous and homogeneous catalysis. The production methods of primary amines are then explored, ranging from the traditional methods to the most recent examples, as well as the merits and drawbacks in each. In addition, production methods of a few non-primary amines are discussed, as they show promise for future research. There will be extensive discussion on N- alkylation of amines with alcohols via hydrogen borrowing technique, as it is more relevant to this study.

2.1. Ammonia, The Nitrogen Source

By definition, amines are a group of chemical compounds with the common feature of possessing nitrogen atoms that are sp3 hybridised with three single bonds to other elements. The simplest amine is ammonia, whereas higher classes of amines include primary, secondary and tertiary amines, with alkyl or aromatic groups replacing one, two or three hydrogen atoms respectively, as shown in Scheme 2.1.

Scheme 2.1. (From left to right) Ammonia, and primary, secondary and tertiary amines, where R = alkyl or aromatic groups

Among these different classes of amines, primary amines are particularly useful as intermediates due to their highly basic nature and ability to establish X-N bonds to form other compounds where X can be C, O or even metal atoms. However, achieving high selectivity of the primary amine can be difficult, because its high basicity can lead to side reactions, which includes the formation of secondary amines and tertiary amines.

The industrial production of amines may involve catalysis, which needs to be conducted efficiently and economically under safe conditions [28]. Ideally, this should involve the use of cheap and abundant metals as catalysts in a continuous flow process, for long- term economical and sustainable benefits, as suggested by the American Chemical

Society Green Chemistry Institute Pharmaceutical Roundtable [29].

Ammonia is perhaps one of the most important bulk chemicals utilised by society, used in the production of valuable nitrogen-containing compounds. The large-scale production of ammonia was pioneered by the German scientist, Fritz Haber (1868 –

1934) in the early 20th century. The Haber Process is an artificial nitrogen-fixing method, whereby nitrogen (N2) from the air and hydrogen (H2) is converted into ammonia (NH3). Before this, nations such as Germany relied on naturally occurring nitrates for fertilisers and explosives from South America, but demand for nitrates increased with a growing global population and the threat of international militancy

[30]. In 1908, Haber first developed an iron-catalysed method of synthesising ammonia from its elements, and in 1914, Carl Bosch, a German chemical engineer converted this method into an industrial process [31]. This process is shown in Scheme 2.2.

Scheme 2.2. Reaction scheme of the Haber-Bosch process with typical industrial conditions

This process is known as the Haber process or the Haber-Bosch process, and was an important development in war efforts. This is because pre-existing methods, such as the

Birkeland-Eyde and the Frank-Caro processes were highly energy inefficient [32-33]. [32][33]

The current global annual production of ammonia is estimated to be 140 Mt, 80% of which is used as fertilisers, further synthesised to urea, ammonium nitrates and

32 phosphates or as anhydrous ammonia for direct application [34]. Ammonia is also used to produce explosives, plastics, synthetic fibres and resins, as well as intermediates for dyes and pharmaceuticals [35]. It can also be directly applied as an excellent refrigerant.

Due to the many applications of this chemical, the physical properties of ammonia are well studied and understood. The phase diagram of ammonia and its solubility curve in toluene are shown in the appendix (Appendix S1).

Properties of ammonia

Ammonia is a colourless gas at room temperature (25 °C) with a strong pungent smell.

It is highly soluble in water in room conditions and forms a basic solution of 880 kg/m3 when saturated. It has a melting point of -74 °C and a boiling point of -33 °C at 1 atm.

The ammonia molecule has the geometry of a trigonal pyramid as predicted by the

Valence Shell Electron Pair Repulsion theory (VSEPR theory), with a bond angle of

106.7° between the N-H bonds (Scheme 2.3). This is slightly smaller than the common tetrahedral bond angle of 109.5°, which is caused by a lone electron pair on the nitrogen atom, which repels the shared electron pairs, resulting in a decrease in the bond angles

[36]. This shape gives the ammonia molecule a dipole moment, allowing the formation of hydrogen bonds which explains ammonia’s high miscibility with water [37].

Scheme 2.3. Molecular structure of ammonia. The lone electron pair repels the shared electron pairs, causing a slight decrease in bond angles.

Ammonia commonly acts as a nucleophile in organic reactions. Primary amines can be formed by reaction with alkyl halides, which will lead to secondary and tertiary amines formations, as the primary amines are generally more nucleophilic than ammonia.

However, catalysis that involves ammonia is challenging in laboratory conditions due to several reasons:

• high toxicity compounded by the volatility of ammonia,

• strong coordination affinities of ammonia to transition metals forming strong

bonds, resulting in highly stable compounds and thus catalytically inert

compounds,

• N-H bond is very strong (ca. 386 kJ/mol) and is difficult to break [38],

• its properties as a reactant are difficult to utilise in a laboratory setting, since

▪ it is highly soluble in water, and

▪ it is a non-ideal gas which makes quantification difficult.

To avoid the direct usage of ammonia, one could employ ammonia substitutes, such as urea, thiourea, triphenylmethylamine (tritylamine), magnesium nitride (Mg3N2) or

lithium amide (LiNH2) [39-40]. [39, 40] However, these surrogates generally suffer from drawbacks such as low molecular efficiency or they produce metallic wastes.

Consequently, anhydrous ammonia is the most atom efficient option.

Summary

In this section, the importance of ammonia has been established, as well as properties of ammonia and the challenges when using ammonia in a laboratory setting. Overall, ammonia is a very well-studied compound due to its utility in many different industries.

Therefore, it is convenient to use ammonia as the amine source in this study.

In the next section, homogeneous and heterogeneous catalysis will be discussed, and their merits and drawbacks will be explored in detail. This would help in discussing the pros and cons when comparing different primary amine production methods.

2.2. Heterogeneous and Homogeneous Catalysis

There are two main types of catalysts, heterogeneous and homogeneous. In heterogeneous catalytic reaction, the catalyst is in the same phase as the reactants. On the other hand, in homogeneous catalytic reaction, the catalyst is in the same phase as the reactants. In this section, both catalyses will be explained, and several examples will be illustrated. Finally, the future of catalysis will be discussed.

2.2.1. Heterogeneous Catalysis

A typical example of a heterogeneous reaction would be a solid catalyst with reactants in their liquid or gas phases. In the typical examples, most catalysis go through the following stages:

• One or more reactants are adsorbed on to the surface of the catalyst at active

sites.

• Interactions between the surface of the catalyst and the reactant molecules

causes more reactive molecules

• This might involve actual reactions between the reactants with the surface, or

some weakening of the bonds in the attached molecules.

• The reaction then happens. The reactants might still be attached to the surfaces,

or one might be attached and hit by the other one moving freely in the fluid.

• The product molecules are then desorbed, where they break away from the

catalyst surface. This leaves the active sites available for the next set of reactants

to attach and react.

A good heterogeneous catalyst must be able to adsorb the react molecules strongly enough for them to react, but not too strongly that the product molecules stay attached to the surface.

Metals like platinum and nickel are good catalysts because they adsorb strongly enough to hold and activate the reactants, but does not hold the products so strongly that they

35 cannot break away. On the other hand, silver is not a great catalyst because it does not form strong enough attachments with reactant molecules; and tungsten isn’t good because it adsorbs too strongly to the reactants.

Example of heterogeneous catalysis

Hydrogenation/Reduction

The simplest example of a heterogeneous catalysis is the reaction between ethene and hydrogen in the presence of a nickel catalyst[41]. In reality, this reaction is rather wasteful because the useful ethene is converted into a relatively useless ethane. This reaction is shown below as Scheme 2.4.

Scheme 2.4. Ethene hydrogenation with a nickel catalyst at 150 ⁰C.

However, this reduction reaction can happen with any alkene containing compound. For example, the hydrogenation of unsaturated vegetable oils to make margarine is conducted at an industrial scale annually[42]. The ethene molecules are adsorbed the nickel catalyst surface as the double carbon bond breaks. Hydrogen molecules break to atoms and are bonded onto the nickel surface. These hydrogen atoms can move around on the surface of the nickel.

If a hydrogen atom diffuses near one of the bonded carbons, the bond between the Ni-

C bond breaks as the C-H bond forms, releasing one end of the original ethene molecule.

When the other Ni-C bond breaks and forms a C-H bond, the product ethane molecule can break free. This space can be used by another reactant molecule to undergo the entire process again.

Another example of heterogeneous catalysis are catalytic converters. These converters change poisonous molecules such as carbon monoxide (CO) and various nitrogen oxides

(NOx) into less harmful gases such as carbon dioxide (CO2) and nitrogen (N2). They are necessary in many automobiles and uses expensive metals such as platinum (Pt), palladium (Pd) and Rhodium (Rh) as the heterogeneous catalyst [43]. A typical catalytic reaction between CO and NO is shown as Scheme 2.5.

Scheme 2.5. Catalytic conversions of poisonous gas molecules to less harmful ones.

The expensive metals are deposited onto a porous ceramic material which maximises the surface area while minimises the amount of metal used. However, catalytic converters are prone to catalyst poisoning, where certain mixture components or side products gets adsorbed onto the catalyst surface strongly, thus preventing further reactions from happening at those sites. Lead is a common catalyst poison for catalytic converters. It coats the porous catalyst structure which stops the catalytic reactions.

Even though lead enabled engines to use higher compression ratios that made cars more powerful, gasoline containing lead additives must not be used in modern vehicles that contains catalytic converters [44].

Heterogeneous Catalyst Deactivation

Heterogeneous catalysts can deactivate under different mechanisms when their active sites lose their functionality via different pathways. They can be classified into five distinct types:

2.2.1 poisoning,

2.2.2 fluid compound formation, fluid-solid or solid-solid reactions accompanied

by transport,

2.2.3 coking or carbon deposition,

2.2.4 mechanical failure, and

2.2.5 thermal degradation.

These deactivation mechanisms can be grouped into three types, where 2.2.1 and 2.2.2 are chemical in nature, 2.2.3 and 2.2.4 are mechanical and 2.2.5 is thermal in nature.

2.2.1. Poisoning

Poisoning is the strong chemisorption of reactants, products or impurities on catalytic sites where they would otherwise be accessible for reactions. The rate of poisoning depends on the concentration of the poison, and the reversibility depends on the strength of the poison adsorption. The poisoning mechanism is a complicated process that involves some of or all the following:

a) physical blockage of the catalytic sites by the strongly adsorbed poison,

b) electronic modification of the nearest neighbouring atoms or beyond by the

poison,

c) restructuring of the adsorbent surface, or

d) hindering surface diffusion of the adsorbed reactants.

Common poisons, such as coke, sulphur and arsenic compounds are strongly and irreversibly adsorbed. Therefore, poisoning is best prevented through purification of the reactant stream with scrubbers, grading or guard bed [45].

2.2.2. Fluid compound formation, fluid-solid or solid-solid reactions

Side reactions in the fluid phase with the catalytic surface may cause:

a) inactive bulk or surface covering compounds (not poisons as not strongly

adsorbed), or

b) leaching, or volatile catalytic compounds that is transported out of the

catalyst.

Other reactions that could happen include catalytic solid support or catalytic solid- promoter reactions, and solid-state transformations of the catalytic phases during the reaction. Examples include the oxidation of Co metal supported by silica due to product water, to form Co surface silicates during Fischer-Tropsch (FT) synthesis at high conversion [46]; loss of Pt by formation of PtO2 during ammonia oxidation on Pt-Rh gauze catalysts [47]; formation of inactive KAlO2 during ammonia synthesis at the

Fe/K/Al2O3 catalyst surface[48]; and reductive transformation of Mo18O52 to Mo4O11 during partial oxidation of propene or acrolein [49]. These forms of chemical catalyst deactivation can be prevented by carful control of reaction conditions, as well as the correct design of catalyst.

2.2.3 Coking or carbon deposition

Fouling is the physical deposition of species from the fluid phase onto the catalytic surface, which results in catalyst deactivation due to the blockages of active sites and/or pores [50]. Fouling can happen in two ways, coking or carbon deposition, and can extend to disintegration of catalyst particles and plugging of the reactor voids, if left unattended[51]. Examples include:

a) large amounts of carbon deposition due to CO disproportionation during

operation at high temperatures and/or at low ratios of H2/CO or steam/C in FT

synthesis, methanation and steam reforming of methane[52], and

b) multilayer accumulation of coke in catalytic cracking on zeolites or

hydrotreating on CoMo/Al2O3 catalysts.

These processes happen differently mechanistically: CO or carbon dissociates on the metal surface to polymerise to undesirable forms such as graphite or carbon filaments; while the coke formation happens by free radical carboncation reactions acid sites, including dehydrogenation, oligomerisation or cyclisation.

To prevent carbon formation, the following steps can be implemented:

a) operation under conditions that minimise formation, e.g. at sufficiently high

H2/CO ratios;

b) optimisation of catalyst design, e.g. optimise zeolite acidity to minimise coke

formation; or

c) purification of the reaction feed to remove precursors that may accelerate

carbon or coke formation, e.g. removal of polynuclear aromatics from the feed

of a hydrocracking or hydrotreating process

Catalyst deactivated by this method can usually be regenerated by low temperature combustion in air.

2.2.4 Mechanical failure

Mechanical failure of catalysts can be caused in several ways [53]:

a) crushing of granular, pellet or monolithic catalysts [54],

b) size reduction and/or breakup of catalyst pallets or granules to produce a fine

powder, especially in fluid or slurry beds [55], and

c) erosion of catalyst particles or monolith coatings at high fluid velocities.

In addition to the loss of catalyst active sites, the mechanical failure of catalysts pallets damage or fracture affects the pressure drop, local temperature and overall efficiency of the process, and would lead to the shut-down of the overall process and subsequent catalyst replacements. Therefore users can use damage models to predict the effect of process parameters, such as packing efficiency, reactor wall roughness and temperature profiles.

On the other hand, users can prevent this from happening by strengthening the catalysts, by means of:

a) adding binders to improve the strength and toughness of the catalyst,

b) increasing catalyst agglomerate strength by advanced preparation methods,

such as sol-gel granulation, spray drying or carefully controlled precipitation,

c) coating of catalyst aggregates with a very strong but porous material, such as

ZrO2, and

d) chemical or thermal tempering of catalyst agglomerates to introduce

compressive stress, which can increase strength and attrition resistance.

2.2.5 Thermal degradation

Thermally induced deactivation of catalysts may result from the following methods:

a) loss of catalytic surface area by crystallite growth,

b) loss of support area due to support collapse, or catalytic surface area due to pore

collapse [56], and/or

c) chemical transformation of catalytic phases to non-catalytic phases [57].

The first two methods are commonly known as sintering, which takes place at high reaction temperatures (> 500 °C) and are generally accelerated by the presence of water vapour. To prevent sintering, a lower operating temperature is recommended.

In the next section, homogeneous catalysis will be explored with its pros and cons demonstrated using examples.

2.2.2. Homogeneous Catalysis

Homogeneous catalysis is where the catalyst is in the same phase as the reactants. This can be in gas, liquid or even solid state. There are several advantages to use homogeneous catalysts over the heterogeneous counter parts [58]:

• the active sites are fully exposed to the reactants, allowing free interaction

between them without any adsorption/desorption pathway,

• this allows mechanistic insights in the reaction and are crucial in understanding

and further developing multifaceted metal mediated transformation [59],

• the homogeneous complexes can then be tuned electronically and sterically by

varying the metal and/or ligands, allowing product selectivity, and

• for exothermic reactions, the heat can be released more readily in a solution,

when compared with heterogeneous catalysts.

However, the inherent disadvantages of homogeneous catalysis are [58]:

• catalyst separation from the product is very difficult for recycle/reuse, which is

highly important per the principles of green chemistry, and

• the temperature of the process is limited by the volatility of the solution and

reagents

Examples of homogeneous catalysis

Polymerisation of 1-alkenes by the Ziegler-Natta catalyst

In 1963, the Nobel Prize in Chemistry was awards to Karl Ziegler for his discovery of the first examples of titanium-based catalysts, and Giulio for using them to prepare stereoregular polymers from propylene. These polymers represent the largest-volume commodity plastics and chemicals in the world. Although this reaction was first discovered with heterogeneous catalysts, homogeneous catalysts were developed over the years.

Metallocene catalysts are organometallic catalysts with derivatives of the organic ligand cyclopentadienyl (Cp). Typically, these catalysts have the composition Cp2MCl2, where

M = Ti, Zr or Hf. The cocatalyst with metallocenes are typically methylaluminoxane

(MAO). These catalysts and all alkylaluminium cocatalysts are unstable in air, and are therefore always prepared and handled under an inert atmosphere.

As mentioned previously, one advantage of homogeneous catalysis is the understanding of the mechanism. This can be demonstrated with the Ziegler-Nata polymerisation, as shown below in Scheme 2.6.

Scheme 26. Ziegler-Nata polymerisation using a homogeneous Ti catalyst.

The simplified metallocene complex, Cp2TiCl2 represents a typical precatalyst, which is unreactive towards alkenes. The complex is activated by MAO to form a

+ metallocenium ion Cp2Ti Me, which allows insertion reactions of C=C bonds of 1- alkene, growing a polymer. The termination can happen via several pathways:

Scheme 2.7. (Top) Ziegler-Nata polymerisation termination with the β-elimination from the polymer chain. (Bottom) Ziegler-Nata polymerisation termination with the β- hydrogen elimination reaction.

Oxidation of ethylene to acetaldehyde with the Wacker process

This process uses palladium(II) chloride as the catalyst to oxidise ethylene to acetaldehyde, and was one of the first homogeneous catalysis with organopalladium chemistry used on an industrial scale (Scheme 2.8) [60].

Scheme 2.8. General scheme of the Wacker process.

The general reaction mechanism of the Wacker process has been debated, but a modern formulation is shown in Scheme 2.9 [61].

Scheme 2.9. Modern formulation of the catalytic cycle of the Wacker process.

It is necessary for copper(II) chloride (CuCl2) to act as an oxidising agent, or Pd(0) metal would precipitate and stop the reaction. Air or other reagents can oxidise CuCl back to

CuCl2 to continue the catalytic cycle.

As mentioned before, the leaching of catalytic species into the liquid phase might be responsible for the actual reactivity of the catalyst. Catalyst leaching is a particularly interesting deactivation mechanism for Pd catalysed C – C, C – O and C – N coupling reactions [62].

In 2010, Akira Suzuki was awarded the Nobel Prize in Chemistry for his work on the

Suzuki-Miyaura reaction, which is a highly active C-C coupling reaction. A huge number of papers and patents had since then claimed the highly effective reaction with

Pd as a catalyst [63-65]. The major shortcoming of many reported systems is the homogeneous nature of the organo-Pd complexes or colloidal Pd. Thus, complete removal from the reaction mixture to avoid metal carryover is intricate and costly, reducing the potential for industrial and commercial implementation conducted in the liquid phase, especially in the field of pharmaceutical synthesis.

Since then, research had been done on the assessment of Pd catalyst leaching in continuous flow [66, 67]. In these works, the catalytic activity of the heterogeneous catalyst was monitored over time, enabled by flow reactors.

2.2.3. Summary and The Future of Catalysis

In this section, heterogeneous and homogeneous catalyses were explained in detail with many examples. This is a constantly expanding research topic as there is great potential to improve on the current state.

Lanzafame et al. [68] suggests the following to be the grand challenges for catalysis:

1) catalysis to address the evolving energy and chemical scenario,

2) catalysis for a cleaner and more sustainable future, and

3) addressing catalysis complexity, which can be divided into:

a) advanced design of novel catalysts,

b) understanding catalysts from the molecular to the material scale, and

c) expanding catalysis concepts

To achieve catalysis for a cleaner and sustainable future, it was suggested to improve the sustainability of chemical processes, in terms of atom economy and improved processes to produce the main intermediates and chemical products/monomers; this area includes:

a) moving towards 100% selectivity;

b) catalysts in novel process design for resource and energy efficiency;

c) novel catalytic processes to reduce eco-impact or risk of fine and specialty

chemical production; and

d) catalysis for novel polymers.

A portion of these are endeavours to be fulfilled in this project. The following section illustrates different primary amine production methods.

2.3. Primary Amine Production Methods

As previously discussed in Chapter 1, Section 1.1, primary alkyl amines are very important in our daily lives and are highly valuable. As such, many production methods were developed over the years. In the follow sections, each of these methods will be explored and discussed in detail:

2.3.1. Alkyl Halide Amination,

2.3.2. Olefin Hydroamination,

2.3.3. Reductive Amination,

2.3.4. Hydroaminomethylation,

2.3.5. Nitrile Hydrogenation, and

2.3.6. Amine Alkylation with Alcohol via Hydrogen Borrowing Technique

A summary of these methods is shown in Scheme 2.10.

Scheme 2.10. Summary of primary alkyl amine synthesis methods, where common examples of reactants are included.

2.3.1. Alkyl Halide Substitution

The most common method to synthesise amines through alkyl halides is reacting alcohols directly with HCl to produce the alkyl halide, followed by halide substitution by an amine, which is usually more varied in method, such as using organic salts

(phthalimide) or metal to catalyse the amination (Scheme 2.11).

Scheme 2.11. General reaction scheme for alkyl alcohol to a halide, and subsequently substitution by an amine to form an alkyl amine.

Gabriel synthesis was one of the first recorded methods of amine synthesis. It begins with alcohol, which traditionally underwent nucleophilic substitution by halides (Cl or

Br) under acidic conditinos, followed by a condensation reaction using potassium phthalimide with amines [13] (Scheme 2.12). The coupled phthalimide is then treated either with hydrazine (Route 1) or acid (Route 2). Both routes yielded the product primary amine, and the by-products were phthalhydrazide hydrazine and benzene-1,2- dicarboxylic acid, respectively.

Scheme 2.12. Gabriel synthesis. An alcohol was converted to the halide and reacted with phthalimide to yield a primary amine.

The Gabriel synthesis was highly selective towards primary amines as higher amines do not react with phthalimide. However, the atom economy was generally poor, as a stoichiometric quantity of phthalyl by-products was generated. Since then, efforts were

48 made to improve this process, including the addition of co-catalysts or development of new reagents [69]. Another example used 3,4-diphenylmaleic anhydride as a catalytic

Gabriel reagent, where no phthalyl waste was produced [70] (Scheme 2.13).

Scheme 2.13. 3,4-diphenylmaleic anhydride as a catalytic Gabriel reagent in an amine substitution reaction.

Of late, tosylates, epoxides and esters have been used in place of alkyl halides in an effort to reduce the use of halides and undesirable salt products. Yadav et al. [69] used tosylhydrazones as the alkylating species instead of conventional halides and CuI as a co-catalyst for carbene insertion into the N-H bond of phthalimide. The process is highly selective towards the primary amine, but required 2 – 4 equiv. of base (Scheme 2.14).

Scheme 2.14. Proposed mechanistic pathway of the Cu-catalysed insertion of carbene into the N-H bond of phthalimide for the amine substitution reaction.

Metal-Catalysed Alkyl Halide Amination

Transition metals can be used to catalyse halide substitution reactions with amines.

Well-known examples include the Pd-catalysed Buchwald-Hartwig cross-coupling and

Cu catalysed Cham-Lam coupling reactions. However, most of these examples involve direct synthesis to yield secondary or tertiary amines [71]. Up to the time of writing, there has been only one example of metal catalysed Buchwald-Hartwig type cross- coupling reaction, which yielded 53 – 99% primary aryl amines as the product [72]

(Scheme 2.15). Although the reaction conditions were rather mild at 50 – 100 °C, a stoichiometric amount of base was used and metal salts were generated as waste products. Metal salts are particularly bad for the environment as they are generally highly soluble in water and are prone to contamination.

Scheme 2.15. Example of aryl primary amine production by Buchwald-Hartwig cross- coupling using a homogeneous Fe catalyst [72].

Chaterjee and Goswami [73] reported a metal-free amination with

[bis(trifluoroacetoxy)iodo]benzene and N-bromosuccinimide to activate O- methylhydroxylamine as an ammonia surrogate to aminate alkyl boronic acids [73].

Although respecTable yields of 65 – 78% were achieved at room temperature, 2 equivalents of the iodo-compound and imide were required for the reaction, which is not atom efficient (Scheme 2.16).

Scheme 2.16. Metal-free amination of alkyl boronic acids [73]. the scope is limited to aryl compounds only, with limited examples in primary amines and no instances of having produced primary alkyl amines.

To summarise, alkyl halide substitution is a promising technique for the synthesis of primary amines. The Gabriel synthesis and its derivations are highly selective to primary amines, but are generally atom inefficient in that stoichiometric amounts of reagents are required. Transition metal catalysed substitution is a highly efficient technique in the cross-coupling of aryl halides or tosylates with amines, but the scope is currently limited as there is currently only one example that produced primary aryl amine as the final product. The metal-free amination of alkyl boronic acids instead of alkyl halides is also demonstrated.

2.3.2. Olefin Hydroamination

Olefin (or alkene) hydroamination represents the addition of R2N-H across C=C bonds and is atom-economical, as no byproduct is produced (Scheme. 2.17). There is a wide scope of amine substrates, including ammonia, primary and secondary aliphatic and aromatic amines, azoles, hydrazines and N-protected amines [74]. This process can be catalysed by transition metals, rare earth metals or alkali metals, and selectivities can be achieved with the functionalisation of unsaturated substrates and catalysts.

Scheme 2.17. A general scheme of hydroamination reaction between ammonia and an alkene.

Early examples of hydroamination with ammonia were found with heterogeneous catalysts. One of the earliest examples was reported by Howk et al. in 1954. Alkenes and ammonia were reacted with sodium under extremely harsh conditions of 175 – 200

°C at 800 – 1000 atm [75]. They achieved low yield at 12 – 34% to the primary amine, with secondary and tertiary amines were the by-products (Scheme 2.18). Therefore, this process is not selective.

Scheme 2.18. An early example of hydroamination using sodium as a reactant [75].

An example of an industrialised process is the hydroamination of isobutylene to tert- butylamine under supercritical conditions, using aluminosilicate catalyst. (Scheme

2.19) This process was commercialised by BASF, with an announced capacity of 8000 mt per annum [76]. This initial patent described primarily aluminosilicate and borosilicate catalysts with pentasil structures. The process conditions were

NH3/isobutylene = 1.5, 300 °C and 300 bar, and the conversions were 9 – 17% with

52 selectivities up to 95%. Subsequently, several patents were described in 1998 - 1999 that used different aluminosilicate based catalysts [51-56]. However, the scope of alkenes remains narrow because isobutylene provides the most stable carbonium ion intermediate; whereas ethylene and propylene form far less stable primary and

secondary carbonium ions, leading to much lower amine product yields. [77][78][79][80][81][82]

Scheme 2.19. Industrialised hydroamination of isobutylene to t-butylamine [78].

More recently, Bertrand et al. reported the first example of homogeneous catalysis for the hydroamination of alkynes and allenes with ammonia using gold (Au) [83]. This example requires a specifically tailored catalyst to overcome the high activation barrier due to the electrostatic repulsion between the lone electron lone pair at the nitrogen atom and the double or triple bond of the olefin (Scheme 2.20). It was reported that by using

4.3 mol% of catalyst and NH3/allene = 40 for 16 hours at 175 °C, 96% conversion was achieved with a primary amine selectivity at 86%.

Scheme 2.20. Homogeneous catalytic hydroamination of allyl groups with ammonia using a homogeneous Au catalyst [83].

Continuing with the use of noble metals, a Pd/Ir dual metal tandem catalyst system was found to be active in catalysing the one-pot, two-step hydroamination of olefins with an ammonia source to give branched primary amines in high yields with only 1 mol% of

Pd and 1 mol% of Ir. This is the first report of a formal one-pot intermolecular hydroamination of olefins to obtain primary amines [84] (Scheme 2.21).

Scheme 2.21. Formal one-pot, two-step hydroamination of olefins using a homogeneous Pd/Ir dual metal tandem catalyst system [84].

The two-step process was found to be necessary to prevent the oxidation catalyst (Pd complex) from being poisoned due to the presence of ammonia. This system achieved moderate to good yields primary amines at 36 – 94% at a moderate temperature of 70

°C.

More recent developments of olefin hydroamination used biocatalysts [59-61]. [85][86][87]

One such example used phenylalnine ammonia lyases (PAL) in tandem with L-amino acid deaminases (LAADs) to produce chiral amines with high selectivity [86] (Scheme

2.22). A high NH3 excess was used (40 equivalents) and the example R groups were electron withdrawing functional groups, which activated the target carbon. Conversion was high at 62 – 80%, and high selectivity to the primary amine was achieved with 98

– 99% ee.

Scheme 2.22. Enantioselective hydroamination of olefins by biocatalysts to a carboxylic acid [86].

In summary, while heterogeneous hydroamination afforded low yields of the target product, homogeneous hydroamination required the use of precious metals with the drawback of difficult catalyst recovery. Perhaps a future development of heterogeneous precious metal catalyst can solve both problems at once. Alternatively, biocatalysts have recently discovered to perform the reaction with high efficiency. However, enzymes suffer from the drawbacks of requiring highly specific biological materials as well as rate limited by the low deformation temperature. In addition, the substrate scope was limited to producing amino acids, and it is unknown whether the synthesis would be successful if the carboxylic acid group was not present.

2.3.3. Hydroaminomethylation

Hydroaminomethylation was first developed by Reppe et al. at BASF in 1949 using stoichiometric amounts of Fe(CO)5 [88]. It describes the process where alkene, amine and syngas (CO and H2) undergo cascade hydroformylation, condensation and reduction sequence, to form a primary alkyl amine with one extra carbon than the starting alkene

[89] (Scheme 2.23). This process is atom-efficient with water being the only by-product.

There are many examples of selective tertiary or secondary amines synthesis, due to the high reactivity between the product amine and aldehyde intermediate. As a result, very few examples were reported for primary amines [16].

Scheme 2.23. A general scheme of hydromethylation of olefins to produce primary amines. Note the increase in the length of the compound by one carbon [89].

A dual metal hydroaminomethylation catalysis with ammonia was described by

Zimmermann et al. in 1999, which was the first example of a selective process towards primary amines with this type of reaction [90]. Rh and Ir were used with the addition of

TPPTS or BINAS ligands (Scheme 2.24). They proposed that [Rh] was responsible for the first hydroformylation step, followed by a fast uncatalysed condensation step, with

[Ir] responsible for the final reduction step. The reaction conditions were 130 °C and 60

– 78 bar. The scope of alkenes in their studies was narrow where only propene, 1-butene and 1-pentene were used. High conversions were achieved at 75 – 90%, and 76 – 87% selectivity to primary amines was achieved.

In 2007, Klein et al reported a catalytic hydroaminomethylation with 1-octene under supercritical ammonia, using a homogeneous bimetallic Rh-Ir catalyst system [91]. The system operated at 210 bar and 140 °C, with oct-1-ene:[IrCl(COD)]2 :[Rh(acac)(CO)2] in a 1000:5:1 ratio and 100 equivalents of acetic acid. It achieved 80% conversion with a selectivity of 60 % towards the primary amine RNH2, where R = alkyl C9 in 16 hours, while limiting the side reactions to only 16% secondary amine. However, a much higher pressure is required to maintain supercritical ammonia compared to other examples.

More recently, a hydroaminomethylation with ammonia to primary amine was reported by Behr et al.[92]. This is particularly interesting as the alkene of choice was limonene, a renewable terpene feedstock. Using a high Rh catalyst loading of 25 mol%, the maximum primary amine yield was low at 25% in a biphasic solvent system. Mild conditions were used, at 130 °C and only 60 bar of syngas without the need for supercritical ammonia (Scheme 2.25).

Scheme 2.25. Hydroaminomethylation of limonene by Behr et al. [92].

In summary, hydroaminomethylation is an atom efficient process in producing primary amines selectively while increasing the compound length by one carbon. However, there are several drawbacks to hydroaminomethylation with ammonia. The reaction requires the use of expensive metals such as Rh or Ir for high yields. Moreover, these catalysts are homogeneous which makes catalyst recycling more difficult. Excess ammonia must be used to ensure the selectivity to the primary amine, which in turn can block catalytic sites and minimise activity, especially for the hydrogenation step. Moreover, the resulting aldehyde intermediates are prone to aldol or other condensation reactions, yielding large amounts of unwanted by-product [91].

2.3.4. Nitrile or Nitro-group Hydrogenation

There are several well-known hydrogenation methods, e.g., use of excess LiAlH4 [93], borohydride reduction of a nitrilium salt [94], use of metal hydride reagents, or perhaps the most direct, high-pressure hydrogenation with a metal catalyst, as demonstrated in

Scheme 2.26 [95].

Scheme 2.26. A general scheme of nitrile reduction with common metal catalysts and

H2 as the reductant.

However, some of these methods have intrinsic shortcomings including stoichiometric metallic waste and hazardous reaction conditions. Working with hydrogen at high pressures and temperatures may be dangerous because it has a wide flammability range, high burning rate, low ignition energy and a non-luminous flame which aggravates the combustion hazards [96]. In addition, the methods may be incompatible with other functional groups such as nitro or allyl groups, and would lead to unwanted side products. As a result, more sustainable and selective methods have been developed recently to overcome this issue.

The high-pressure hydrogenation process is well understood and can be applied to nitriles to produce amines. Catalysts such as Ni [97], Pt, Fe, Cu, Co and Rh are usually used under high pressures of hydrogen, with the reagents also in the gas phase [98].

Ideally, the nitrile should be reduced to selectively produce primary amines, but the amines can undergo further reactions and form higher amines, as shown in Scheme 2.27.

This a major drawback as it lowers the selectivity towards the primary amine.

Scheme 2.27. Potential side reactions of nitrile reduction.

Homogeneous Hydrogenation Catalysis with Base Metals

In 2016, Lange et al. demonstrated a Fe pincer catalyst that can achieve highly selective reduction of nitriles to the primary amine [99]. A high H2 pressure of 30 bar was used with a moderate temperature of 70 °C and yields of 44 – 95% were achieved (Scheme

2.28).

Scheme 2.28. Nitrile reduction with a homogeneous Fe complex by Lange et al. [99].

More recently, Adam et al. achieved selective hydrogenation of nitriles to primary amines using a cobalt catalyst with tri-dentate phosphine ligands [100]. High activity was found to be achieved with tris[2-(dicyclohexylphosphino)ethyl]phosphine (Scheme

2.29). Various aromatic and aliphatic nitriles were successfully reduced to primary amines, with 100% conversion and 61 – 99% yield.

Scheme 2.29. Selective reduction of nitriles using a cobalt phosphine catalyst by Adam et al. [100].

Heterogeneous hydrogenation catalysis

On an industrial scale, heterogeneous catalysis is preferred to produce amines from nitriles. The following examples are novel catalytic processes that demonstrated highly selective hydrogenation of nitriles to primary amines.

In 2016, Chen et al. developed a stable Co/α-Al2O3 catalyst for highly selective hydrogenations of nitrile and carbonyl functional groups [101]. A wide range of aromatic and aliphatic nitriles were reduced to primary amines at a high temperature of

130 °C under 40 bar NH3 for 2 hr in the presence of aqueous ammonia, and high yields of 90 – 98% were achieved. By-products included the secondary amine as well as the intermediate imine. The catalyst was found to remain active after 8 recycling experiments, which is great for the longevity of the batch process.

In 2017, Kobayashi et al. reported selective hydrogenation of nitriles to primary amine salts catalysed by a supported Pd catalyst under continuous flow conditions [102]

(Scheme 2.30). A low hydrogen pressure of 0.5 bar was used in their flow reactor, which was heated to 60 °C. High yields of the ammonium salt were achieved at 97 – 100%.

The catalyst was able to remain active for more than 300 hours (TON > 10 000) without loss of selectivity.

Scheme 2.30. Flow scheme of selective hydrogenation of nitriles to primary amines catalysed by a supported Pd catalyst, reported by Saito et al. [102].

In the following example, a non-traditional reductant was used for the reduction of the nitrile to primary amines. Gandhamsetty et al. [103] reported a boron-catalysed silylative reduction of nitriles to produce amines and imines, shown in Scheme 2.31.

B(C6F5)3 was used as the catalyst (1-3 mol%) with hydrosilanes (2.5 equiv.) as the reductant. Although the boron catalyst is highly efficient, high stoichiometry of silanes was required. Besides, acid workup was required to achieve the ammonium salt to remove the silane groups, which would require another work up to achieve the amine.

These steps would lead to considerable waste and are thus environmentally unfriendly.

Scheme 2.31. Boron-catalysed silylative reduction of nitriles to primary amines [103].

In summary, nitrile reduction generally has fewer side reactions, as only the intermediate imine can undergo side reactions, compared to other reactions where carbonyl intermediates area also involved. Non-metal catalysed processes were demonstrated with good efficiency. In addition, the substrates are not limited to the gas phase, allowing a wider scope for nitriles to be hydrogenated.

2.3.5. Reductive Amination

Carbonyl groups can be converted to amines with an amine source under reductive environment. However, only ammonia can provide primary amines, and this poses the challenge with the primary amine being more reactive than ammonia, leading to over- alkylation. In addition, molecular hydrogen is hazardous, and chemical reductants usually result in high amounts of waste and low atom economy. This process had been performed by both homogeneous and heterogeneous catalysts [104]. An example of reductive amination of an aldehyde using a stoichiometric reductant NaCNBH3 is shown in Scheme 2.32.

Scheme 2.32. Example of reductive amination of an aldehyde to form a primary amine.

Furfural is used as a model aldehyde by a few research groups for the reductive amination process. This is because furfural is a renewable starting material derived from lignocellulosic biomass, which does not compete against those with food or oil-based origins. Also, the product primary amine, furfurylamine (FAM) is widely used in the manufacture of pharmaceutical drugs, pesticides, fibres, perfumes, etc. [105].

Back in 1986, Ayusawa et al. investigated this reaction using Raney Ni in NH3 (liq.) under 87 atm H2, yielding >97% furfurylamine [106]. Using a cheap Ni catalyst, the reaction required harsh conditions to achieve the high yields. Interestingly, recent FAM production examples required milder conditions but involved the use of Ru catalyst.

Using furfural as the model aldehyde, Komanoya et al. screened a wide range of metals on different supports, including Ru, Rh, Pd, Ni, Cu, Ag and Pt supported by Nb2O, SiO2,

TiO2, C, Al2O3, ZrO2 and MgO. They found that Ru/Nb2O3 gave the highest yield to

FAM at 99% under 1 bar NH3 and 40 bar H2 [107]. On the other hand, Ebitani et al. reported a Ru heterogeneous catalyst that was effective in the reductive amination of

63 furfural with aqueous ammonia under hydrogen [108]. They achieved a 60 % yield of

FAM using a PVP-capped 5wt% Ru-supported HAP catalyst in 25 % NH3 (aq) under

2.5 atm H2 gas at 100 °C. However, this catalyst had lower performance with common aromatic aldehydes at 31 – 42% yields. The difference in yields between these two examples is rather large (99% compared to 60%). This could be due to the difference in their catalysts, but may also be due to the difference in reaction conditions, as they used different ammonia sources. Komanoya et al. used gaseous ammonia while Ebitani et al. used the aqueous form, and the excess water could be hindering the forward reaction.

This is an interesting topic which will be discussed in greater detail in the case of alcohol amination via the hydrogen borrowing technique in Chapter 2, Section 2.6.

Later, Delmas et al. reported the synthesis of primary amines by one-pot reductive amination of aldehydes using Zn/HCl, albeit proposing a different pathway (Scheme

2.33) [12].

Scheme 2.33. Potential pathways for reductive hydrogenation.

Route A describes the common pathway, where the aldehyde reacts with ammonia to form the primary imine, and subsequently gets hydrogenated to the primary amine. The problem is that the intermediate imine is reactive and can react with the product amine or the starting alcohol to form secondary amines. Alternatively, route B describes another approach, where the more stable oxime was first synthesised, and subsequently

64 hydrogenated by Zn/HCl to produce the primary amine, eliminating the possibility of secondary amine formation.

This example was performed with furfural as the model aldehyde, giving high yields of

FAM at 99%. Other examples of alkylamines were synthesised at 80 – 99% yields as well. Although the selectivity to primary amines was very high, a major drawback of this method is the stoichiometric reagents required for the hydrogenation step and the associated waste.

So far, all reductive amination that produced primary amines required H2 as a reductant.

Although the following examples do not produce primary amines, they have demonstrated alternative reductants that may potentially be used in future to produce primary amines.

Carbon monoxide (CO) was recently used as a reducing agent for reduction amination.

Using a Rh catalyst, Chusov et al. reported this reaction carried out without the use of an external hydrogen source [109]. They claimed that CO has an advantage over H2 because CO is more readily available in the steam methane reforming process, whereas a water-gas shift reaction is required to produce more H2. In addition, their catalyst performed better with CO compared to H2 (Scheme 2.34). High yields were generally achieved with 56 – 98% yield in 4 hours. Even though there are merits in using CO as a product from the steam methane reforming process, CO is a poisonous gas and is not ideal as a reagent.

Scheme 2.34. Scheme of reductive amination using CO or H2 as reductants.

Formic acid (4 equivalents) was also used as a reductant with Au/TiO2 [110] 6 – 24 hours for the synthesis of secondary amines with 88 – 98% yields at 60 °C. Efforts using non-hydrogen reductants can also be seen, e.g. CO and formic acid. Although these might bring other problems such as toxicity and/or the excessive quantities required, it is a good sign to see alternatives to hydrogen, for more flexibility in reactant variety.

Lastly, a biocatalyst was developed recently for the synthesis of chiral primary amines

[111]. The mechanism is shown in Scheme 2.35. Amine dehydrogenases (AmDHs) are a new class of enzymes to catalyse the reduction amination of ketones and aldehydes with high conversions (72 – 99%) and high selectivity (>99%). Although a coenzyme

(nicotinamide (NAD)) was necessary, the only by-products were water and carbonate.

No molecular hydrogen was used in the reaction, as the reducing agents and nitrogen source originate from the reaction buffer solution of ammonia/ammonium formate.

Curiously, a reductive amination Ni-Al alloy in water without the need of external reductant was reported by Shafer et al. as shown in Scheme 2.36 [112]. Aqueous ammonia was used as the ammonia source as well as the hydrogen source. However, their claim that “a fresh Al surface can readily react with water used as a solvent to yield hydrogen gas (under ultrasound)” is questionable.

Scheme 2.36. Reductive amination with Ni-Al alloy under ultrasound.

Summary

Reductive amination is a selective process in production primary amines. There are many examples that uses homogeneous or heterogeneous catalysis, as well as a stereoselective biocatalyst system. However, H2 is required as the reductant for primary amines, as well as aldehydes as the reagent. This is not ideal as both reactants are usually unstable and care is required when handling them. CO was demonstrated to work as a reductant, but for secondary amine productions in the example given. Overall, reductive amination is an atom-efficient process with a plenty of potential in primary amine production.

2.3.6. Amine Alkylation with Alcohol via Hydrogen Borrowing Cycle

Alcohols are generally readily available in bulk and are potential alkylating candidates.

However, the inert nature of the hydroxyl group poses a challenge as it may not be easily replaced via nucleophilic substitution. Usually, alcohols are first pre-treated and substituted with better leaving groups, such as halides, tosylates and triflates, or converted to other functional groups, such as nitriles or aldehydes, before further reactions as illustrated in the previous sections (Scheme 2.37).

Scheme 2.37. Alcohol preparations before conversion to primary amine.

Taking the atom economy into consideration, these reactions tend to require extra workup, produce more waste and thus are atom inefficient. For a sustainable future, a

“greener” or environmentally friendly process should be the goal [18].

Hydrogen borrowing technique is one such example with high atom efficiency [113].

This technique, also known as the hydrogen auto-transfer process, uses alcohol directly as the starting material without the need for any pretreatment, with water generated as the only by-product. The mechanism of the hydrogen borrowing technique is as follows

(Scheme. 2.38).

Scheme 2.38. Mechanism of the hydrogen borrowing technique [113].

The hydrogen borrowing technique goes via 3 steps

1) the alcohol undergoes oxidation to an aldehyde or ketone, which then

2) reacts with an amine to through a condensation reaction to give an imine, and

3) the imine subsequently undergoes reduction by the initially generated hydrogen to yield the product amine.

The hydrogen is effectively “borrowed” from the starting alcohol and “auto- transferred” to the imine to produce the final amine, thus the name of the mechanism.

There are a few advantages in the hydrogen borrowing technique. Firstly, the reactive aldehydes/ketones or imines are produced in situ, which is safer as carbonyl/iminyl compounds suffer from poor stability (e.g. oxidation of aldehydes to carboxylic acids, or self-enolisation which leads to unwanted side products), which eliminates the need to store the reactive species, making the overall process safer. Secondly, this process does not require additional reductants, such as hydrogen or sacrificial hydrogen donors, which is advantageous since most of these reductants are toxic or flammable and may generate high amounts of waste. Lastly, water is the only byproduct, which is attractive

69 compared to the aforementioned methods, which generated toxic wastes [114]. These advantages are in alignment with the concepts of green chemistry in terms of safer reaction conditions as well as the reduction in waste generation [18].

History of the hydrogen borrowing cycle

One of the first reported examples of amine alkylation with alcohols was in 1932 by

Winans and Adkins [115] (Scheme 2.38). Cyclohexylamine was reacted with ethanol to yield the secondary and tertiary amines in 55% and 17% selectivities respectively. A high excess of hydrogen at 75 atm was used with 12 mol% metal Ni. This was described as “reaction of amines with alcohols with the elimination of water”, though it did not fully describe the mechanism. The authors also suggested that “the same factors seem to govern the activity of nickel for the catalysis of the alkylation of amines as govern its activity for hydrogenation, since a preparation of catalyst which was rather inactive for hydrogenation was also rather inactive for alkylation”, thus suggesting that strong hydrogenation/dehydrogenation catalysts would perform well in these reactions.

Scheme 2.39. One of the first reported examples of alkylation of amines with alcohols by Winans and Adkins [115].

This reaction was later done with homogeneous catalysts by Grigg et al. [116] and

Watanabe et al. [117] using Ru, Rh and Ir, which are known to be active in hydrogenation reactions. Both examples used aniline as the amine source while reacting with C2-4 alcohols, yielding a mix of primary and secondary amines, as shown in Scheme

2.39. Other examples of homogeneous Ru catalysts include Ru(cod)(cot), which achieved a higher selectivity to mono-alkylation, and [Ru(cymene)Cl2]2 with dppf

[bis(diphenylphosphino)-ferrocene], which was successful in regard to N-alkylate

aromatic and aliphatic amines at a lower temperature of 110 °C. (Scheme 2.40)

Scheme 2.40. Early homogeneous reactions using a homogeneous Ru catalyst in the

hydrogen borrowing cycle [117].

In Table 2.1, the recent advances in the catalytic syntheses of primary amines from

alkylation of ammonia using the hydrogen borrowing technique over the past 10 years

are shown, including homogeneous and heterogeneous examples. In this Table, the

catalyst, reaction conditions, conversion and selectivities and turnover number (TON)

are shown. A range of alcohol conversions and selectivities are given where provided.

Specific conversions or yields are also illustrated in terms of specific alcohols, such as

benzyl or 2-octyl alcohol because these can be used for direct comparison, as labelled.

TON was calculated based on the produced amine per catalyst metal used = [produced

amine]/[M].

Table 2.1. Recent examples of catalysis of N-alkylation of ammonia with alcohols

Entry Catalyst Conditions Conversion (Selectivity)a TONb Reference

1 Ru pincer 0.1 mol% cat, reflux, NH3 96 – 100% (69 - 96 %) 1000 [118]

complex (7.5 atm), 13 - 32 hours, 100% (87%), R = benzyl

toluene

2 Ru pincer 1 mol% cat, 140 °C, , 32 - 95% (51 – 100%) 95 [119]

complex NH3/ROH = 35, 21 – 64

hours, 6 mL cyclohexane

3 Cu(OH)x/Al2O3 2 mol% cat, 135 °C, 80% yield (2° amines), 40 [120]

NH3/ROH = 0.25, 6 hours, 1 R = benzyl

atm Ar, 2 mL mesitylene

4 Ru(OH)3.TiO2 0.6 mol% cat, 141 °C, 87% yield (3° amines), 145 [121]

NH3/ROH = 0.2, 6 hours, 1 R = benzyl

atm Ar, 0.5 mL mesitylene

5 Ni/Al2O3 1 -5 mol% cat, 140 – 160 °C, 93 – 100% (77 – 96%) 87 [122]

NH3/ROH = 2.2, 24 – 72 99% (78%), R = benzyl

hours, 4 g o-xylene

6 Ni/CaSiO3 1 - 5 mol% cat, 140 - 170 °C, 70 – 88% yield 88 [123]

NH3/ROH = 2.2, 20 hours, 4 70% yield, R = benzyl

g o-xylene

7 NiCuFeOx 92 mol% cat, reflux, 60 –90% yield (2° amines), 0.91 [124]

NH3/ROH = 1.2, 24 hours, 2 R = benzyl & alkyl

mL xylene (NH4)2CO3

8 Ni/Al2O3 55 mg cat, 160 °C, 69% (90%), R = 1-octyl 87 [125]

NH3/ROH = 5.8, 4 hours, 3

mL o-xylene, 600 rpm

9 2NiCe/Al2O3 2.3 mol% cat, 180 °C, 82% (78%), R = 1-octyl 69 [17]

NH3/ROH = 5.8, 24 hours, 3

mL o-xylene, 600 rpm

10 ADH/AmDH 5 mol% cat, 30 °C, 85 – 95% (>99%) 19 [126]

NH3/ROH = 20, 48 hours,

0.5 mL CO2NH4 buffer

(1.005 M)

a R = alkyl, selectivity to primary amines, unless specified otherwise. b TON was

calculated by [amine]/[M]

In 2008, Gunanathan and Milstein developed a homogeneous acridine-based pincer complex [RuHCl(A-iPr-PNP)(CO)] for the N-alkylation of ammonia and alkyl alcohols to synthesise primary amines (Table 2.1, Entry 1, Scheme 2.41). A very low catalyst loading of 0.1 mol% was used to achieve a high conversion at 100% for benzyl alcohol and 87% selectivity to the primary amine. [118] The side-products were the corresponding secondary amines.

Scheme 2.41. Structure of acridine-based pincer complex [RuHCl(A-iPr-PNP)(CO)] used in the alkylation of ammonia via the hydrogen borrowing cycle [118].

Gunanathan and Milstein performed more N-alkylation of ammonia with alcohols in the presence of water and achieved a slightly higher selectivity at 95% with slightly longer reaction time (18 hours vs 12 hours). They attributed this improvement to the on-water effect, where water played a part in catalysing the reaction. This is interesting because the addition of water should shift the thermodynamic equilibrium backwards, where less secondary amine would form, which may also explain the increase in selectivity to the primary amine.

In 2010, Pingen et al. [119] used a homogeneous Ru catalyst with a different pincer ligand to selectively synthesise primary amines via the hydrogen borrowing technique, using liquid NH3 at a large excess (35 equiv.) (Table 2.1, Entry 2, Scheme 2.42). This batch process was carried out in an autoclave, and conversion of 32 – 95% was achieved for secondary alkyl alcohols, and selectivity to the primary amines ranged from 51 –

100%.

Scheme 2.42. Structure of pincer ligand used by Pingen et al. in the homogeneous Ru catalytic system in the alkylation of ammonia via the hydrogen borrowing cycle. [119]

Although moderate to high TON could be achieved with homogeneous Ru catalysts

(TON = 95 – 1000), they suffer from difficulty in catalyst recovery which is complicated with the use of expensive noble metal catalysts [127]. The following examples are focused on heterogeneous base metal catalysts.

In 2010, Mizuno et al. [120] used a heterogeneous Cu catalyst for the N-alkylation of ammonia with alcohols in mesitylene (Table 2.1, Entry 3). A high yield of 80% towards the secondary amines was achieved at an NH3/ROH ratio of 0.25. Aqueous ammonia was used as the ammonia source, and it was not reported that water affected the reaction.

Theoretically, water inhibits the condensation reaction in the hydrogen borrowing cycle.

The potential adverse effects of water in the hydrogen borrowing cycle will be discussed in more details in Chapter 5, section 5.6 (Effect of Water). A modest TON of 40 was achieved in the synthesis of secondary amines.

In the same year, Mizuno et al. [121] demonstrated a hydrogen borrowing N-alkylation of ammonia with alcohols using a supported ruthenium hydroxide catalyst

(Ru(OH)x/Al2O3) under anaerobic conditions. An excess of ROH was used in these reactions (ROH/NH3 = 5), which promoted the formation of higher amines. As a result, primary amines were not observed, and tertiary amines were produced at 87% (Table

2.1, Entry 4) [128]. It should be noted that, under aerobic conditions, oxidative synthesis of nitriles directly from alcohols or aldehydes was achieved with NH3 using the same catalyst [129].

In 2013, Shimizu et al. [122] developed a heterogeneous Ni/Al2O3 catalyst that achieved high selectivity to primary amines using gaseous ammonia. High conversions were achieved at 93 – 100% as well as high selectivity to primary amines at 77 – 96% with 1 mol% catalyst loading and in 13 – 29 hours (Table 2.1, Entry 5). In particular, a conversion of 99% and selectivity of 78% were achieved using benzyl alcohol, albeit at a higher catalyst loading of 5 mol% and longer reaction time of 72 hours. This catalyst was reusable after three repeats with a slight decrease in performance (conversion down from 96% to 86%). Gaseous ammonia was used in the experiment at an NH3/ROH ratio of 2.2, which led to the higher selectivity to primary amines.

In the following year, Shimizu et al. [123] developed a Ni catalyst supported by CaSiO3 with high selectivity to primary amines, again using gaseous ammonia. They achieved a 70 – 88% yield to primary amines using a NH3/ROH ratio of 2.2; in particular, 70% yield on benzyl amine with a higher molecular loading (5 mol%). This catalyst was also effective in reacting aryl and alkyl amines with alcohols.

These two examples by Shimizu et al. [94, 95] had a better performance with secondary alkyl alcohols compared with benzyl alcohols. This is apparent from the need for higher catalyst loadings (5 mol% compared to 1 mol%) and/or longer reaction times (72 hours compared to 24 hours), and usually a lower conversion. This will be highlighted in

Chapters 5 and 6 as their catalysts’ performances differed from our catalysts.

A trimetallic catalyst was used by Cui et al. [124] to achieve the hydrogen borrowing

N-alkylation of ammonia with alcohols, namely NiCuFeOx (Table 2.1, Entry 7). The weight ratio of Ni:Cu:Fe species was about 3:1:1 and a high catalyst loading of 92 mol% was used in the reaction with NH3 and ROH. It should be noted that urea was used as an NH3 surrogate in these reactions. Using a low NH3/ROH of 1.2, they achieved yields of 60 – 90% for aryl and alkyl alcohols to the secondary amines. They suggested that the synergism between the Ni, Cu and Fe species might be crucial to realising the

75 borrowing hydrogen transformation. A low TON of 0.92 was achieved with this system due to the high catalyst loading.

In 2017, Tomer et al. [17] reported a Ni/Al2O3 catalyst with similar performances as that of Shimizu et al. [94]. Ni/Al2O3 with a range of different Ni loadings (2 – 15%) was synthesised and tested, and 10 wt% was found to be the most selective towards primary amines. Their model alcohol was n-octanol and they achieved a conversion of 69% and

90% selectivity to the primary amine (Table 2.1, Entry 8).

More recently, Tomer et al. [17] developed a bimetallic NiCe/Al2O3 catalyst which was also selective towards primary amines (Table 2.1, Entry 9). As in the previous example, they tested catalysts with a range of different Ni loadings (0.5 – 8 wt%) and found that the 2 wt% Ni catalyst was the most effective. n-Octanol was the model alcohol and 32% conversion and 79% selectivity to the primary amine were achieved.

There were recent examples of biocatalysts using the hydrogen borrowing technique, albeit with a slightly different catalytic cycle, as shown in Scheme 2.43. Turner et al.

[126] used a two-enzyme cascade system, namely an alcohol dehydrogenase (ADH) operating in tandem with an amine dehydrogenase (AmDH). At a high ammonia excess

(NH3/ROH = 20) using an aqueous ammonium buffer, this system achieved good conversion at 85 – 95%, with excellent selectivity at >99% (Table 2.1, Entry 10). The products were also enantiopure, so that the process is deemed highly valuable for pharmaceuticals and biomedicine. A modest TON of 19 was achieved using this system.

Scheme 2.43. Two-enzyme cascade regime for the hydrogen borrowing alkylation of

+ amines. Conditions. 30 C, 24 – 48 hours, 5 mol% enzymes, 2M NH4 /NH3 [130].

Although the turnover numbers (TON) were calculated and shown in the Table for direct comparison, it should be noted that they were calculated with respect to the amount of bulk catalyst (in mol), as opposed to the number of active sites. For instance, the TON of homogeneous Ru pincer complex was very high at 1000 (Table 2.1, Entry 1), compared to 145 of heterogeneous Ru catalyst (Table 2.1, Entry 4).

There is a recurring challenge towards high selectivity to primary amine when reacting with ammonia. When R is an electron donating group, which is generally the case, the more nucleophilic primary amine product proceeds to react with the starting alcohols, producing unwanted secondary amines (Scheme 2.44). Depending on the steric effect of the R group, the secondary amine may proceed to react and form tertiary amines.

Scheme 2.44. Nucleophilicities of ammonia, a primary amine and a secondary amine in increasing order.

To summarise, selectivity towards primary amine is highly dependent on the NH3/ROH excess according to literature. Selectivity to primary amines was achieved at NH3/ROH

> 2.2, and selectivities to tertiary and secondary amines were achieved at NH3/ROH ≤

1.2. This will be revisited in Chapter 5, Section 5.2 (Thermodynamics).

2.3.7. Critical Comparison of the Amine Production Methods

As demonstrated in the Sections 2.3.1 to 2.3.6, there are many methods to synthesise

primary amines, and some recent examples were discussed. However, most of them

suffer from drawbacks such as hazardous conditions, atom inefficiency and/or unwanted

side reactions, such as the use of high-pressure hydrogen in hydroaminomethylation and

reductive amination. To ensure a sustainable future in chemistry, the use of

heterogeneous catalysts should be encouraged, as well as the elimination of

stoichiometric reagents and high atom efficiency. Furthermore, Lanzafame et al. [68]

suggested that developing 100% selective process to produce main intermediates and

chemical products/monomers is important to tackle challenges in catalysis for a

sustainable future.

Table 2.2 illustrates some of the features of these processes for a critical comparison.

These features include the type of catalysts used (homogeneous, heterogeneous or

biocatalyst).

Table 2.2. Comparison of the six primary amine production methods.

Entry Reaction name Homogeneous catalysis Heterogeneous catalysis Biocatalyst

1 Alkyl halide substitution Transition metals Transition metals No

2 Olefin hydroamination Noble metals Aluminosilicate Yes

3 Hydroaminomethylation Noble metals None No

Nitrile or nitro-group Transition and Noble metals 4 Transition metals No hydrogenation and boron

5 Reductive amination Transition and Noble metals Noble metals Yes

NH3 alkylation with 6 Noble metals Transition and Noble metals No alcohol

1) Traditional alkyl halide substitution is highly selective to primary amine, but

require stoichiometric amounts of reagents and produces much waste.

Transition metal catalysed reactions are more atom-efficient, but its scope is

currently limited.

2) Olefin hydroamination can be achieved using heterogeneous or homogeneous

catalysis. Homogeneous catalysis afforded high conversions and selectivities to

primary amine, but requires expensive precious metals with difficulties in

recovery; heterogeneous catalysis currently affords low yields to the primary

amine using aluminosilicate.

3) Hydroaminomethylation adds one additional carbon to the product amine,

which might not be desirable if a small product molecule is required. The

current method uses homogeneous catalysis with noble metals only, which

makes it a very costly process.

4) Nitrile or nitro-group hydrogenation requires a reductant, which is typically

flammable and dangerous to the working environment. This process is also non-

selective and can potentially reduce other unsaturated bonds in the substrate.

5) Reductive amination also requires a reductant. There is still room for

improvement in heterogeneous catalysis, as they are only available using

expensive noble metals.

6) N-Alkylation of ammonia with alcohol is available with transition metals for

both homogeneous and heterogeneous catalysis. Although this reaction is well

explored with larger amines (e.g. aniline, dimethylamine) as the amine sources,

ammonia is used relatively less with limited examples. The heterogeneous

catalysis gave high selectivity and appeared to be controlled by NH3/ROH

molar ratios, but are generally slow as they require a long reaction time.

2.4. Conclusion

In this chapter the following have been discussed:

• ammonia as a very important source of nitrogen, and is the most atom efficient

form of amine as a reactant,

• a general overview of heterogeneous and homogeneous catalysis and some

examples, as well as their advantages and disadvantages, with heterogenous

catalysis being more favoured for the recyclability of the catalyst, which agrees

with the practices of sustainable chemistry, and

• the many different methods to produce primary alkyl amines and a critical

comparison between them It was identified that the N-alkylation of ammonia

with alcohols is a reaction with great potential, as it can be:

o performed with heterogeneous catalysts,

o performed using flow chemistry,

o highly atom efficient,

o selective towards primary amines.

Chapter 3: Aims and Objectives

3.1. Aims

This project aims to develop a new process to produce primary amines. The process should include the following aspects:

• safe working environment, and

• atom efficiency

The literature review in the previous chapter illustrated the potential of hydrogen borrowing technique compared with the other chemical processes in the synthesis of primary amines. The hydrogen borrowing technique is more atom efficient with water being produced as the only by-product.

Scheme 3.1. Mechanism of hydrogen borrowing technique

Water is produced as a by-product in the condensation step of the hydrogen borrowing cycle. In a batch reaction, water is accumulated in the system and causes a slower rate of reaction because the accumulation of water discourages the formation of imine. The use of a plug flow reactor will reduce the accumulation of water by-product and amine products that can inhibit the forward reaction, thus faster reaction rates and fewer by-

81 products. In addition, the reaction parameters such as temperature and pressure can be accurately measured and controlled. These features allow the reactions to be performed more effectively at shorter residence times compared to a batch process, making it a more effective operating process for scaling up and increasing the production capacity.

It is also easier to test for long term catalyst performance under flow conditions. The use of hydrogen borrowing technique in a flow reactor can be an elegant solution to fulfil the aims illustrated above.

Many catalysts, including Ru and Ni, promote the N-alkylation of NH3 by alcohols, generally with Ni performing at a lower conversion. However, by carefully controlling the starting conditions such as temperature and reaction stoichiometry, which strongly influences the reaction outcome, high conversions of alcohol and highly selective monoalkylation of NH3 may be achieved.

3.2. Objectives

The aim of the project will be achieved by accomplishing the following objectives:

(i) Understanding the thermodynamics of the system,

The thermodynamics of the reaction system should be studied to understand the physical state of the reaction mixture under reaction conditions. This is important to ensure that the reaction mixture is in liquid phase, because a gas-liquid mixture would require a different reaction mechanism. The product distribution will also be calculated at reaction completion, and the theoretical results can be compared with the experimental results. These theoretical calculations can be achieved using simulation software such as ASPEN-plus and CAPE-OPEN flowsheeting environment (COFE).

(ii) Finding a suitable catalyst,

As discussed in the literature review (Chapter 2, Section 2.6), noble metal catalysts such as Au and Ru are more selective towards secondary and tertiary amines, where base

82 metal catalysts such as Ni and Cu are more selective towards primary amine. The specific tasks of this objective will involve screening commercial catalysts, synthesising new catalysts and modifying existing catalysts with the aim to selectively synthesise primary amine. The model batch reaction will be the N-alkylation of ammonia, NH3 with benzyl alcohol, BnOH.

(iii) Studying reaction kinetics to control rate and selectivity,

By understanding the kinetics of the reaction system, the conversion of alcohol and selectivity to amines can be controlled by adjusting the reaction conditions such as temperature, reactant stoichiometry and residence time. The reaction mechanism for the

N-alkylation of NH3 with BnOH can be established by analysing the batch reaction mixture over time and determining the reaction rate constants using the experimental results.

(iv) Adapting, designing and constructing a continuous flow system that

allows safe handling of ammonia and achieving higher yields than the

batch system.

A bench-top continuous flow system is adapted from a previous work, and designed and constructed for the N-alkylation of NH3 with alcohols, which must deliver NH3 safely.

The model reaction is the N-alkylation of NH3 with BnOH with initial reactions conditions established in batch reactions and refined to adapt for flow conditions. The catalyst activity is monitored over extended periods of time to observe potential catalyst deactivation.

Chapter 4: Experimental Protocol and Methodology

4.1. Introduction

In this chapter, all materials used in the project are first illustrated, along with their suppliers. Experimental methods, including batch and flow reactions and catalyst preparations, computational and analytical methods used in this project are explained, which will be referred to in later chapters and sections.

4.2. Materials

Unless stated otherwise, all chemicals used were of standard reagent grade and used without further purification. Table 4.1 lists all chemicals and material used in this project and their corresponding suppliers.

Table 4.1: List of chemicals and materials used in the project Chemicals and materials Supplier

Au/TiO2, 1 wt% Strem Chemicals, U.K.

Ni/Al2O3/SiO2, 65 wt% Sigma Aldrich

Ru/Al2O3, 1 wt% Sigma Aldrich

Pt/Al2O3, 1 wt% Sigma Aldrich

HIFUEL Ni/Al2O3, 16 wt% Ni Fisher Scientific, U.K.

NiMo/AlxOy Millán et al. [131]

Bentonite clay, (montmorillonite, Al2O3.4SiO2.H2O) Fisher Scientific, U.K.

Methyl cellulose Sigma Aldrich

Benzyl alcohol, anhydrous, 99.8% Sigma Aldrich

Benzaldehyde, ≥99% Sigma Aldrich

Benzyl amine, 99% Sigma Aldrich

Dibenzyl amine, 97% Sigma Aldrich

Tribenzyl amine, >99% Sigma Aldrich

2-Picolyl alcohol (2-pyridinemethanol), 98% Sigma Aldrich

3-Picolyl alcohol (3-pyridinemethanol), 98% Sigma Aldrich

2-Methoxyl benzyl alcohol, 99% Sigma Aldrich

o-Xylene, 97.8% Fisher Scientific, U.K.

3-Methoxyl benzyl alcohol, 98% Sigma Aldrich

2-Phenylethanol, ≥98% Sigma Aldrich

2-Octanol, ≥97% Sigma Aldrich

Nitrogen, compressed, 100% BOC

Argon, compressed, 100% BOC

Ammonia, anhydrous, 100% BOC

Ammonia hydroxide, aqueous, 28 wt% Sigma Aldrich

Nickel Nitrate Hexahydrate, 99.999% Sigma Aldrich

Sodium silicate (aq), 1 g/L Sigma Aldrich

Ni (1000 mg/L) ICP standard solution, TraceCERT Sigma Aldrich

Acetic acid, glacial, ≥99.85% Sigma Aldrich

Hydrochloric acid, 37% Fisher Scientific, U.K.

Nitric acid, 65% Fisher Scientific, U.K.

Methanol, 99% Sigma Aldrich

Ethanol, absolute, 99.98% Sigma Aldrich

H2 was supplied by an in-house H2 generator (Precision Hydrogen 200)

4.3. Experimental

4.3.1. Batch Reaction

The reactions were typically performed in a 30 mL batch reactor with a sample port.

(Figure 4.1) The sample port consisted of a 1/16” tube which reached the solution inside the reaction. The inlet of the sample port had a filter (25 μm mesh) to separate catalyst from the sample solution, which otherwise would cause blockages in the 1/16” sampling tube. The total internal volume of the reactor was 42 mL, but typically only 30 mL of solution was used for reactions, leaving 12 mL of headspace. There were two different batch reaction setups used in this work using different NH3 sources: aqueous or anhydrous.

Sample dispenser

Figure 4.1. Batch reactor combined with the condensation tube.

There are pros and cons for each setup: the use of aqueous NH3 allowed safety in handling and transporting, as well as ease of quantifying NH3 by weight or volume.

However, aqueous NH3 intrinsically possesses 72% w/w H2O, which has adverse effects in the substitution step of the hydrogen borrowing cycle. Anhydrous (or pure) ammonia

86 is more difficult to handle and quantify, as it is gaseous under standard room conditions, compounded by its toxicity when uncontained. However, these challenges were overcome by a condensation method, as detailed later in the chapter.

Typical batch reaction with aqueous NH3

In the reaction that involved aqueous NH3, the commercial 65 wt% Ni/Al2O3/SiO2 catalyst (60 mg Catalyst 1), stirrer, 30 mL of 0.2 M BnOH solution in o-xylene (6.0 mmol) and aqueous NH3 solution (105 µL) were placed in the Teflon liner. The reactor was then closed and purged three times with N2 to remove air. The reactor was then sealed at 4.2 bar under N2, which was then heated at 160°C, stirring at 1000 rpm. The pressure was typically 18 bar at 160 °C. Samples were collected over time through the sample port. On completion, the reactor was cooled to room temperature before pressure release.

Typical batch reaction with anhydrous NH3

In the reaction that involved aqueous NH3, the catalyst extrudites (100 mg, 36%wt Ni), stirrer and 30 mL of 0.2 M BnOH solution in o-xylene (6.0 mmol) were placed in the

Teflon liner. The reactor was then closed and purged three times with N2, followed by purging three times with NH3 to remove air. The reactor was sealed and NH3 (l) was introduced into the reactor with a condenser tube (Figure. 4.2). This was an SS316 ¼” tubing (0.12 m) cooled under acetone/ice bath (-10 °C) while connected to NH3 for 20 minutes through a valve. The amount of condensed NH3 was determined by the weight difference (0.70 g, 41 mmol). The NH3 was introduced into the batch reactor aided by a heat gun (150 °C). The condenser tube was then detached before the reactor was sealed and heated to 160°C while stirring at 1000 rpm. The pressure was typically 12 bar during reaction. Samples were collected over time through the sample port. On completion, the reactor was cooled to room temperature before pressure release.

L = 0.12 m Submerged in acetone/ice bath when condensing

Figure 4.2. Condensation tube to introduce anhydrous NH3 into the batch reactor.

Repeated batch reactions

The commercial 65 wt% Ni/Al2O3/SiO2 catalyst (Catalyst 1) was reused to verify if it remained active for multiple runs. The catalyst was separated from the cooled reaction mixture by filtration, subsequently washed with ca. 30 mL of acetone and 30 mL of DI water, then dried under vacuo for 20 mins. The catalyst was then reused in subsequent reactions without further treatment.

Blank Batch Reactions with Homogeneous Ni (II) species

A method to test whether the reaction is catalysed by the heterogeneous or homogeneous catalyst is as follows:

If the leached catalyst is detected in a solution from a heterogeneous catalyst, use the solution to dissolve reactants to test for reactivity with the absence of the heterogeneous catalyst. If any reaction happened, it would mean that the homogeneous catalyst was also responsible for the overall reaction. The solvent can also undergo elemental analysis (e.g. Induced Coupled Plasma) to detect if any metal content was present.

In some batch reactions with the commercial catalyst, the solution appeared light green after the reactions, suggesting that homogeneous Ni might have leached out. However, the Ni concentration of the sample was lower than the Induced Coupled Plasma (ICP) detection limit of 0.2 ppm. Blank reactions were performed to test if homogeneous

88 species contributed to the reaction. This was done in two ways, the use of Ni(NO3)2 and the use of reaction solvent.

The reactions were performed in a 30 mL batch reactor with a sample port.

Ni(NO3)2·6(H2O) (110 mg, 10 mol%Ni), stirrer, 30 mL of 0.2 M BnOH solution in o- xylene (6.0 mmol) and NH3 solution (105 µL) were placed in the Teflon liner. The reactor was then closed and purged three times with N2 to remove air. The reactor was then sealed at 4.2 bar under N2 and then heated at 160°C, stirring at 1000 rpm. The pressure was typically 18 bar at 160°C. On completion, the reactor was cooled to room temperature before pressure release. However, no reaction occurred.

The blank reactions that used reaction solvent were performed as follows. Catalyst extrudites were heated and stirred in o-xylene for 72 hours in the sealed reactor.

Afterwards, the catalyst was removed from the solvent, BnOH was added into the solvent and heated for an additional 72 hours with NH3 in the reactor for 72 hours.

Although the solution appeared green, no conversion was observed. This may suggest that Ni might have leached out into the liquid phase as Ni(II) is green. This was observed by Zheng et al. [132], where they discovered green solutions when Ni is in their ethanol water systems. Additionally, ICP for Ni was carried out with the solution, but it was not detected, which suggests that concentration of the leached Ni content was lower than the detection threshold (<0.2 ppm).

4.3.2. Flow Reaction

The flow reactor used in this project was a custom-made flow reactor and its scheme is shown below in Figure 4.3 The components of the flow reactor and their description are shown in Table 4.1 and a labelled picture of the flow reactor is shown in Figure 4.4. The justification of each components and choices will be discussed in greater detail in chapter 6.

Figure 4.3. Schematic of the customised flow reactor

Table 4.1: List of components used in the customised flow reactor and their description.

Components Description Supplier

BnOH by Sigma 0.1 - 0.6 M BnOH dissolved in BnOH in o-xylene Aldrich, o-xylene by degassed o-xylene Fisher Scientific, U.K.

Two Gilson 305 HPLC pumps, pump

HPLC pumps head sizes 5 cc & 25 cc for NH3 (l) and Gilson

BnOH solution respectively

Connected to NH3 (g), 20 mL metal

cylinder immersed in: Condenser Swagelok a) N2 (l) during condensation;

b) Water (25 °C) during warm-up

Temperature Range = 25 – 750 °C Omega controller

Furnace and heating blocks to

Heater maximise heat conduction to the Customised in-house

reactor

Reactor 0.28 m 1/4’’ SS316 tubing Swagelok

Thermocouples (Type K) were used to External temperature monitor heater and reactor Omega monitors temperatures

20 μm filter to prevent particulates Filter Swagelok from being flushed downstream

Pressure relief valve Pressure relief at 80 bar Swagelok

Pressure gauge Range = 0 – 200 bar Omega

Back pressure Range = 0 – 175 bar Swagelok regulator

Gas outlet connected to acid trap,

before ventilation. The liquid fractions Gas-liquid separator Swagelok were collected in 1 mL centrifuge

sample holders or 250 mL beakers

Acid trap 1 M HCl aq sln Fisher Scientific, U.K.

Pressure gauge Solutions/solvent s

Gas-liquid separator

HPLC Temperature pumps controller Back pressure regulator

Pressure relief valve

Acid trap 1 M HCl aq sln

Figure 4.4. Labelled photograph of the flow reactor

Experimental

Benzyl alcohol solutions (0.1 – 0.6 M) were prepared with degassed o-xylene. Ammonia was pre-condensed in a condenser (20 mL) at 77 K for 20 minutes and allowed to warm up for 2 hours to room temperature prior to reaction. Each charge of condensation accounted to ca. 10 mL of liq. NH3. Two HPLC pumps were used to deliver the alcohol solution and liq. NH3 respectively, with the pressure controlled by a back-pressure regulator (Swagelok K-series). A pressure relief valve was installed between the reactor and the back pressure regulator and set at a relief pressure of 80 bar. A gas–liquid separator was used before the outlet to remove NH3 from the liquid phase at ambient pressure. The exhaust gas was treated by an acid scrubber (0.1 M HCl solution, 1.0 L) before being vented away, while the sample solution was collected in a flask.

The internal voidage of the packed bed was calculated by the mass difference between a packed reactor when dry and when filled with solvent: a voidage of 78% was calculated (solvent volume of 2.8 mL). The total internal volume of the entire system

(from pump to sample port) was 40 mL.

Flow reactions: Catalyst reduction procedure

2.0 g of catalyst extrudites (Catalyst 2) was loaded into a clean reactor SS-316 tube

(0.28 m), supported by glass wool plugs. The tube was vertically mounted in a heating block. The inlet of this reactor was connected to the reduction gas supply, while the outlet was connected to ventilation through SS316 tubing. The reactor was heated at

500 °C under a reduction gas flow of 10% H2/N2 gas mixture at 40 mL/min for 2 hours.

The gas flows were controlled by mass flow controllers. The reactor was cooled to room temperature before reconfiguring to allow liquid flow. The flow scheme for this prereduction procedure is shown in Figure 4.5.

Figure 4.5. Schematic of the flow reactor configuration for catalyst reduction

To prepare for liquid flow, the reactor was reconfigured. The H2/N2 gas supply and the ventilation outlet tubing was replaced by the solution feed inlet and liquid outlet respectively, as shown in Figure 4.3.

Pump calibrations

The Gilson 305 HPLC pumps were calibrated to make sure that the liquids were being pumped at an accurate rate. The flow rate of the 25 cc pump head was measured over

0.05 to 1 mL/min at 60 bar with o-xylene, and the 5 cc pump head was measured from

0.001 to 0.1 mL/min at room pressure using an inverted 1 mL syringe with water. Water was used because it is very difficult to measure the volume of liquid ammonia without specialised tools. The calibration curves are shown below as Figure. 4.6.

Figure 4.6. Calibration curves for Gilson 305 HPLC pumps, 25 cc (left) and 5 cc pump heads

In both cases, the displayed flow rates were slightly higher than the actual flow rates, by a factor of 1.07 and 1.09 for the 25 cc and 5 cc pump heads respectively. In this study, the reported flow rates were adjusted for this mechanical error.

The compressibility of water vs ammonia

The Z values for water and ammonia are 0.230 and 0.242 respectively [133]. The liquid compressibility factors are similar, and it is assumed that they behave similarly when transported by HPLC pumps.

Residence time calculations

The residence time (τ) can be calculated according to eq. 4.1.

푉 τ = 푑 (eq. 4.1) 푣

Where Vd is the dead volume within a packed cartridge, and v is the applied flow rate.

The voidage (휀) is the ratio of the dead volume (푉푑) and the total inner volume (푉푡표푡) of the reactor, as shown in eq. 4.2.

푉 휀 = 푑 (eq. 4.2) 푉푡표푡

Combining the first two equations, we can get the equation for residence time as

휀×푉 τ = 푑 (eq. 4.3) 푣

Total internal volume (Vtot) = 0.22 m x area of ID of 1/4” SS tubing:

1/4” OD = 0.25”, Wall thickness = 0.035

ID = (0.25 – 0.035)”x2 = 0.18” = 0.00457 m

0.00457 Internal volume = 22 × 휋( )2 = 3.61 × 10−6 푚3 2

Then the packed bed was weighed before and after wetted with o-xylene to find out the filled volume in the packed bed:

Wetted packed bed in tube = 0.053624 kg

Dry packed bed in tube = 0.051160 kg

O-xylene mass = 0.002464 kg

O-xylene density = 880 kg/m3

Void volume = 2.8 × 10−6 푚 = 2.8 ml

2.8×10−6 푚3 Therefore, voidage = = 78%, and τ = 5.6 mins at a flow rate of 0.5 3.61×10−6 푚3 ml/min.

Flow Reactor Characterisation

Reynolds number

Reynolds number (Re) is a dimensionless number that gives a measure of the ratio between inertial and viscous forces under flow conditions. It describes whether the conditions lead to laminar or turbulent flow. The equation below assumes a fluid through a packed bed of approximately spherical particles of diameter D [134].

휌푣 퐷 푅푒 = 푠 (eq. 4.4) 휇 where:

휌 = Fluid density = 761 kg/m3 vs = superficial velocity = 1.27x10-3 m/s

D = diameter = 0.004 m

휇 = dynamic viscocity = 0.00625 kg/m s (from COFE simulation)

Re = 2.48 (dimentionless)

At an Re value of 2.48, which is < 10, it describes laminar flow under the reaction conditions. This is important as it ensures a plug flow model where no back mixing is assumed.

Flow reactor profile: dispersion curve

A series of dispersion profiles were carried out using the flow reactor, at different flow

rates. Dispersion profiles are important to understanding when the solvent or solution

will fill the reactor, which can help plan reactions.

Pure o-xylene was flown at a high flow rate (5 mL/min) through the reactor to fill it

with o-xylene. Once free of impurities, a BnOH solution of known concentration was

flown through the reactor at a much slower flow rate of 0.5 mL/min. The results were

then analysed with GC-FID. An example is shown below as Figure. 4.7.

% Max concentration

Time (Minutes) Figure 4.7. Dispersion profile of the flow reactor at 0.5 mL/min

4.3.3. Catalyst Preparation and Preliminary Performances

Preparation of Ni catalyst by Coprecipitation

The commercial 65 wt% Ni/Al2O3/SiO2 catalyst (Catalyst 1) demonstrated great conversion and selectivity to the primary amine. However, there was an issue with the small catalyst size. The unmodified form of this catalyst is a fine powder (38 – 90 μm) and would not be appropriate for employment in a packed bed due to the high back pressure as well as potential catalyst runoff, which could lead to reactor shut-down and devastating damage to equipment [135]. In an attempt to replicate this high Ni loading catalyst with larger particle size, an extensive search on the production method of the commercial 65 wt% Ni/Al2O3/SiO2 was carried out, which led to a catalyst synthesis patent by Oudejans et al. [136]. Using this co-precipitation method, catalysts of 43 wt%

Ni/Al2O3/SiO2 and 85 wt% Ni/Al2O3/SiO2 were synthesised. The synthesised catalysts had a larger particle size (diameter ~ 1 mm) compared to the commercial 65 wt%

Ni/Al2O3/SiO2 (< 65 μm) and could easily be eroded to form smaller. The 43 wt% catalyst was green and the 85 wt% catalyst was black after the final calcination procedure. The synthesised Ni/Al2O3/SiO2 catalysts were then ground down, sieved and separated to sizes of 38 – 90 μm and 90 – 250 μm. 20 mL of aqueous solutions of

Ni(NO3)2 (35 mg Ni/mL) and Na2CO3 (100 mg/mL) were continuously pumped at equal flow rates (1.0 mL/min) into a vigorously stirred round bottom flask (250 mL) at a temperature of 20 °C, where nickel hydroxide/carbonate was precipitated. The pH of the suspension in this flask was 9.0, controlled by the addition of ammonia solution (7 wt%). The mixture was then heated at 66 °C for 30 min while 80 mL of aqueous solution of Al(NO3)3 (68 mg Al/mL) was added at a rate of 2.6 mL/min, and the pH was kept at

8.4. At this point, the average Al/Ni atomic ratio was 0.15. The suspension was then heated at 97 °C for 30 minutes while 80 mL of Na2SiO3 solution (15 mg Si/mL) was added (2.6 mL/min), during which the pH was kept at 8.9. The average Si/Ni molar ratio at this point was 0.15. The suspension was subjected to an ageing process and was

terminated after 180 min. The suspension was then filtered and the green filter cake was

washed with deionized (DI) water (200 mL). The washed cake was then dried under

vacuo, washed with acetone and dried at room temperature. The catalyst was then

treated under 5% H2/N2 at 400 °C for 30 minutes and used without further treatment.

In this method, the flow rates of solutions were controlled by syringe pumps.

Performances of the Coprecipitation Catalysts

The catalysts were tested in the model NH3 N-alkylation reactions with alcohols with

no further treatment in batch reactors, and the results are shown in Table 4.2.

Table 4.2. NH3 alkylation with BnOH by synthesised and commercial Ni catalysts. Entry Catalyst Particle Sizes BnOH Selectivities (%)

(μm) Conversion (%) BnNH2 Bn2NH

a 1 65 wt% Ni/Al2O3/SiO2 38-90 96 97 3

2 43 wt% Ni/Al2O3/SiO2 90-250 4 100 0

3 85 wt% Ni/Al2O3/SiO2 38-90 25 75 25

4 85 wt% Ni/Al2O3/SiO2 90-250 18 93 7

Conditions: 96 hours, 1000 rpm, 0.60 M BnOH in o-xylene (30 mL), 10 mol% Ni,

a NH3/BnOH = 7, 4.2 bar at 25 °C. Commercial catalyst, 72 hours

The results with the commercial catalyst are shown as a reference point with high

conversion at 96% and high selectivity to BnNH2 at 97% (Table 4.2, Entry 1). 43 wt%

Ni/Al2O3/SiO2 was inactive with a very low conversion was achieved at 4% with a high

selectivity to BnNH2 at 100% (Table 4.2, Entry 2).

On the other hand, 85 wt% Ni/Al2O3/SiO2 showed higher conversions at 25% and 18%

for particle sizes 38 – 90 μm and 90 – 250 μm respectively. (Table 4.2, entries 3 & 4)

The slight increase in conversion with smaller particle size could be due to the higher

surface area. Interestingly, the selectivities to BnNH2 were 75% and 93% respectively.

100

Since the 85 wt% catalyst showed more promising results than the 43 wt% one, the 85

wt% catalyst underwent reduction prior to reaction to improve its activities.

Prereduction procedure (for batch reactions using synthesised catalysts, 85 wt% Ni and 43 wt% Ni)

Catalyst reduction was carried out in a quartz tube (d = 4 mm, l = 0.12 m), where the

catalyst was held by quartz wool, and heated at 500 °C under 10% v/v H2/He flow for 5

hours. The tube was allowed to cool to room temperature under He, then the inside of

the quartz tube was flushed with reaction solution under He and the reduced catalyst

was charged into the reactor.

Performances of the Coprecipitation Catalysts (After reduction)

Table 4.3. NH3 alkylation with BnOH by synthesised (reduced) and commercial Ni

catalysts

Entry Catalyst Particle BnOH Selectivities (%)

Sizes (μm) Conversion (%) BnNH2 BnNH2

a 1 65 wt% Ni/Al2O3/SiO2 38-90 96 97 3

2 85 wt% Ni/Al2O3/SiO2 38-90 25 75 25

3 85 wt% Ni/Al2O3/SiO2 (Reduced) 38-90 20 93 7

4 85 wt% Ni/Al2O3/SiO2 90-250 18 93 7

5 85 wt% Ni/Al2O3/SiO2 (Reduced) 90-250 32 94 6

Conditions: 96 hours, 1000 rpm, 0.60 M BnOH in o-xylene (30 mL), 10 mol% Ni,

a NH3/BnOH = 7, 4.2 bar at 25 °C. Commercial catalyst, 24 hours

The reduced catalysts showed little improvement in activity. The reaction with reduced

38 – 90 μm catalyst saw a slight decrease in conversion from 25% to 20%. (Table 4.3,

Entry 3) This could be due to pacification, where the smaller particle sizes were more

prone to oxidisation when exposed in air, forming catalytically inactive NiO species,

101 which caused the slight decrease in activity. Selectivity to BnNH2 increased from 75% to 93%.

Surprisingly, the reduced 90 – 250 μm catalyst saw a much higher increase from 18% to 32% (Table 4.3, entries 5). This could be because the larger particle size allowed partial pacification of the catalyst on the outer surface, with the inner portion remaining in a reduced form. The reduction activated the outer layer of the catalyst, allowing the

Ni to be observed with X-Ray Diffraction (XRD) (2θ = 46.9°, Figure 4.8). Selectivity to BnNH2 showed a slight increased from 93% to 94%.

NiO 43.6 85 wt% Ni/Al O /SiO 2 3 2 Synthesised, reduced

Ni 46.9Ni 65 wt% Ni/Al O /SiO 2 3 2 Commercial, untreated

Figure 4.8. XRD patterns for Ni catalysts used in batch reactions

It was found that the synthesised catalyst was not as active because there was a large amount of NiO (2θ = 43.6°) compared to Ni (2θ = 46.9°), [137] as determined by XRD analysis, even after the reduction procedure, whereas for the commercial catalyst, Ni was detected in large amounts with very little NiO. Although the reduction method improved the selectivity and conversion of the synthesised catalyst, the improvement was not significant enough to match the activities of the commercial catalyst. In light of this, a different approach was adopted to tackle this problem.

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Catalyst Extrudites

Catalyst Extrudite Production Method

This extrusion method modified from the work of Melero et al. [138], where they agglomerated a zeolite with bentonite clay to form macroscopic structured catalyst particles to be used in continuous epoxidation processes on a fixed bed reactor. The extrudites were made with an active catalyst, an inorganic agglomerant clay and an organic additive as a binder.

In this work, 65 wt% Ni/Al2O3/SiO2 catalyst (2.4 g), bentonite (1.6 g) and methyl cellulose (0.2 g) were carefully mixed together in the dry state under the fume hood. 6.0 mL of 3 wt% acetic acid (aq) was added slowly manually with a syringe while constantly stirring to form a homogeneous paste. The material was then transferred into a 20 mL syringe which was used as an extruder. The extruded catalyst rods were placed onto a clean polythene sheet and air-dried at room temperature for 12 hours, then broken into 1 mm fragments before being dried at 120 °C for a further 2 hours. The extrudites contained 36 %wt Ni as confirmed with ICP. The catalyst extrudites will be known as

Catalyst 2 throughout the project. The activity of Catalyst 2 were tested in batch reactions.

Figure 4.9. The process of making the 36 wt% Ni Catalyst extrudites (Catalyst 2)

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Catalyst particle size control

Catalyst size was controlled by gently grinding the catalyst extrudites (ca. 1.0 g) up with mortar and pestle, then sieving through sieves of specified mesh sizes (250 and 600

µm). This process was repeated several times to acquire the amount needed for the reactions. The sieved particles were then washed with de-ionised (DI) water to remove debris that was electrostatically attached to the particles and dried at 120 °C for 2 hours before further use.

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4.4. Calculations

Conversions were calculated by

[푅푂퐻] × 100% (eq. 4.5) [푅푂퐻]+[푃푟표푑푢푐푡]푇

Selectivity of product 1 was calculated by

[푃푟표푑푢푐푡] 1 × 100% (eq. 4.6) [푃푟표푑푢푐푡]푇

where [Products]T = conc. of products and [ROH] = conc. of alcohol.

Note that for secondary and tertiary amines, the conversions and selectivities were calculated per mole of Bn involved, i.e. Bn2NH would count as 2 equiv. and Bn3N would count as 3 equiv. of benzyl groups.

Turnover Frequency (TOF)

푚푚표푙 푐표푛푐 퐵푛푁퐻 푝푟표푑푢푐푒푑 ( ) 2 푚퐿 ⁄ 푚푚표푙 푁푖 푇푂퐹 = 푚퐿 (eq. 4.7) 푓푙표푤 푟푎푡푒 ( ) 푚푖푛

Turnover number (TON)

푚푚표푙 푐표푛푐 퐵푛푁퐻 푝푟표푑푢푐푒푑 ( ) 2 푚퐿 ⁄ 푚푚표푙 푁푖 푇푂푁 = 푚퐿 ∙ 푟푒푎푐푡𝑖표푛 푡𝑖푚푒 (푚𝑖푛) (eq. 4.8) 푓푙표푤 푟푎푡푒 ( ) 푚푖푛

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4.5. Computational Methods

4.5.1. Phase Diagram Simulations

Two software was used to predict the phase diagrams of the o-xylene and ammonia binary system, namely COFE and ASPEN-plus.

COFE

COFE (the CAPE-OPEN Flowsheet Environment) is a software that translates graphical user interface to a chemical flow-sheeting environment. It has a sequential solution algorithm using and displays properties of streams in a chemical process. This enables the definition of reaction conditions such as the flow rates of components, temperature and pressure, and the software can predict the physical properties of the reaction mixture, such as phase compositions (liquid or vapour), enthalpy and the combined volume using the properties packages built in the system.

Physical properties of NH3 and o-xylene were extracted from the Thermodynamics for

Engineering Applications (TEA) package, an in-built properties package. The properties simulation took place under the following steps and is illustrated in Figure 4.10:

• NH3 was mixed with solvent, o-xylene, in the reactor,

• Stream heated in the heater, then,

• Stream “collected” after the cooler

NH3 Mixed Heated Collection o-xylene Reactor Heater Cooler

Figure 4.10. Simulation flow sheet in COFE with NH3 and o-xylene as the components.

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Table 4.4. Conditions of each stream in the COFE simulation

Reactor Name Function Temperature (°C) Pressure (Bar)

Reactor Mixing 25 1

Heating and Heater 160 1 - 60 pressurising

Cooling and Cooler 25 1 depressurising

The physical conditions were determined as shown in Table 4.4 and the phase compositions were then calculated. Each “stream” shows the calculated physical properties and most important stream after the heater, indicated as heated which shows the state of the reaction mixture during the reaction. These were tested over different stoichiometries by changing the flow rates of NH3 and o-xylene. Properties of pure o- xylene and NH3 (l) were used in the simulation, and the combined flow rate was 0.1 mL/min. The calculated individual flow rates under different stoichiometries are shown in Table 4.6. Properties of pure o-xylene were chosen for this calculation because benzyl alcohol concentration was low (at 0.1 M) and contributions were assumed negligible.

3 Density of NH3 (l) = 609 kg/m [139] was considered when calculating the flow rates.

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Table 4.5. Calculated individual flow rates of the NH3/o-xylene mixture

NH3 flow rate o-xylene flow rate

NH3/BnOH = (mL/min) (mL/min)

1 0.001 0.099

3 0.004 0.096

5 0.007 0.093

7 0.009 0.091

10 0.013 0.087

The physical data, including the vapour and liquid phase compositions, for NH3/BnOH

= 1 - 10 was then collected after the “heater” at the “heated” stream at different pressures

(indicated by arrow). An example of the data representation is shown in Figure 4.11.

Figure 4.11. Physical properties available from the COFE simulation. The “heated” flow stream indicated the mole fractions.

Pure o-xylene was applied in this model as BnOH was assumed to not interfere with the mixture. The conditions were temperature at 160 °C, pressures ranging from 0 to 60 bar and the calculations were carried out over NH3/BnOH molar ratios from 0 to 1. The

108 concentrations of BnOH solution used in the experiment ranged from 0.1 to 0.6 M, and the concentration of 0.6 M was selected in this simulation because it would establish the highest operating pressure at all concentrations. The results for 0.6 M BnOH solution in o-xylene were simulated and are presented below in Figure 4.12.

NH3/BnOH = 1 2 3 4 5 7 10

Figure 4.12. COFE simulated phase diagram of different stoichiometries of NH3/BnOH solutions. Conditions 160 C, pressure = 0 – 60 bar. From left to right: NH3/BnOH = 1,

2, 3, 4, 5, 7, 10.

As shown in the graph, the higher the NH3/BnOH ratio, the greater the pressure required for the mixture to become 100% liquid. For example, NH3/BnOH =1 requires 10 bar for the mixture to become fully liquid, whereas for NH3/BnOH = 10, it requires 60 bar. The thickened line represents NH3/BnOH = 7, which was the operating condition, which shows that at 40 bar onwards, the mixture becomes 100% liquid. This helps confirm that the reaction mixture was in liquid phase at 60 bar, which was the operating condition.

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ASPEN

Background

ASPEN is a process simulation software package widely used in industry today. Given a process design and an appropriate selection of thermodynamic models, ASPEN uses mathematical models to simulate the performance and details of the process. In this project, Aspen plus V8.4 was used to predict the properties of NH3/o-xylene mixture at a fixed pressure (60 bar) over different stoichiometries.

Methods

Many Equation-Of-State (EOS) methods are available for the simulation of a binary mixture. An EOS is a thermodynamic equation that relates state variables, which describes the state of matter under a given set of physical conditions, such as pressure, volume, temperature, and internal energy. The ASPEN has a database for the components containing the state variables, allowing calculations and predictions of the state of the component mixture based on the selected EOS. The descriptions of Ideal gas law and SR-polar are described below. The pure component parameters used in the models were derived from the ASPEN Plus Pure Component Data Bank.

IDEAL gas law

The ideal gas law is an EOS that describes a hypothetical ideal gas. It is a combination of the following laws: Boyle’s law, Charles’s law, Avogadro’s law and Gay-Lussac’s law [140]. It is an empirical law which can also be derived from the microscopic kinetic theory. The equation for the ideal gas law is as follows:

푃푉 = 푛푅푇 (eq. 4.8)

Where P = pressure, V = volume, n = amount of species (in moles), R = the universal gas constant, and T = the temperature in K. However, the assumption is that the

110 components are in gas phases acting as perfect gases, where very small or no interactions occur. However, ammonia acts as a non-ideal gas and a more accurate EOS is required.

Schwartzentruber-Renon-Polar (SR-Polar)

The Schwartzentruber-Renon-Polar (SR-Polar) equation of state is of the Redlich-

Kwong type, which is an empirical equation that relates temperature, pressure and volume of gases. In addition, SR-polar includes a volume translation to reproduce the liquid molar volume more accurately [141].

For a pure component I, the equation of state is defined as:

푅푇 푎 (푇) 푃 = − 푖 (eq. 4.9) 푣−푏푖 (푣+푐푖)(푣−2푐푖+푏푖)

Where

2 1 (푅푇푐푖) 푎푖(푇) = 1⁄3 ∙ 훼푖(푇푟푖) (eq. 4.10) 9(2 −1) 푃푐푖

1⁄3 2 −1 (푅푇푐푖) 푏푖 = ∙ − 푐푖 (eq. 4.11) 3 푃푐푖

1⁄2 훼푖(푇푟푖) =

1⁄2 1 + 푚(휔푖)(1 − 푇푟푖 ) + (1 − 푇푟푖)[푝1푖 + 푝2푖(1 − 푇푟푖)] 푤ℎ푒푛 푇푟푖 ≤ 1 { 1 푑푖 exp [(1 − ) (1 − 푇푟푖 )] 푤ℎ푒푛 푇푟푖 ≥ 1 푑푖

(eq. 4.12)

2 푚(휔푖) = 0.48508 + 1.55171휔푖 − 0.15613휔푖 (eq. 4.13)

푚(휔 ) 푑 = 1 + 푖 + 푝 (eq. 4.14) 푖 2 1푖

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푝1푖and 푝2푖are pure component parameters which are fitted to vapour pressure data

Simulation results

Aspen simulations were carried out to verify the phase composition at the reaction conditions. The SR-Polar EOS was selected on Aspen for the verification of the state of the mixture. The SR-Polar EOSwas derived from the Schwartzentruber–Renon equation, with additional considerations for polar components at high pressures. Other non-ideal methods were tested and produced a similar profile to SR-Polar. The Aspen simulation was carried out for NH3/o-xylene mixture at 60 bar, 50 – 450 °C, and 0 – 1 mole fraction of NH3/BnOH. The resultant plot using SR-Polar is shown below as

Figure 4.13.

A linear region is observed between 0.1 and 0.28 mole fraction (NH3), which is interesting as it suggests that there is no V+L region, but a spontaneous transition from liquid to gas. This behaviour is also observed with other non-ideal mixtures such as methanol and dimethyl carbonate [142]. From 0.28 mole fraction (NH3) onwards, the

112 curves diverge similarly to that of the ideal gas method with three phases, with V+L region separated by the dew point and bubble point curves.

When considering the reaction conditions of 160 °C and 0.02 mole fraction NH3, marked with a red cross on the graph, the reaction mixture was shown to be in the liquid phase using the SR-Polar method.

Both simulations, by COFE and Aspen-plus, led to the conclusion that the reaction mixture was in liquid phase under the reaction conditions, which were 60 bar, 160 °C and 0.02 mole fraction NH3.

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4.5.2. Curve Fittings for Kinetic Studies

Kinetic modelling is an important tool in understanding reaction mechanisms. In this work, the experimental data collected in batch reactions of different conditions (e.g. catalysts, temperature or reactant stoichiometries) was fitted into a defined reaction pathway using a modelling software called Berkeley MadonnaTM, where the reaction rates can be calculated. These reaction kinetic constants can then be compared with each other and the reaction mechanism model can be improved over time.

Berkeley Madonna

Berkeley MadonnaTM is a mathematical modelling software developed at the University of California in Berkeley. It can numerically solve ordinary differential equations and difference equation. It allows for predictions in reaction profiles by applying appropriate rate laws and formulating kinetic equations by manual input of proposed reaction steps.

After applying the reaction details, including the initial conditions and experimental data, values of one or more parameters can be approximated using a curve-fitting feature. The graphical interface can minimise deviations between the model and experimental data.

Simulation Method

The reaction pathway for N-alkylation of NH3 with BnOH is shown in Scheme 4.1. This three reaction model was used as the initial approach to kinetic modelling.

Scheme 4.1. Model of the N-alkylation of NH3 with BnOH with labelled kinetic constants

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The reactions were first defined using the in-built chemical equations function, where the reaction steps and initial component concentrations were defined by the user, and the initial guesses for kf & kr are usually defined as 1.00. (Figure 4.13=4)

Figure 4.14. In-built chemical equations function interface, where the reaction steps and initial concentrations are defined.

For simplicity, components in the reaction mechanism were represented by A – F, as indicated in bold in Scheme 4.1.

In this reaction, BnOH (A), BnNH2 (C), Bn2NH (E) and Bn3N (F) were recorded over time and analysed by gas chromatography-flame ionisation detector (GC-FID). The results were then imported to the software and the curves were fitted according to the predefined reaction pathway, and the kinetic constants for each reaction approximated.

The classical 4th order Runge-Kutta method (RK4) was used as the approximation method. The full details of this method can be found in the appendix (S3). The quality or wellness of a curve fit is determined by the RMS deviation between the data and the best run. This deviation is the root mean square of the differences between individual

115 data points in the dataset and the corresponding points in the run. Therefore, the smaller the RMS, the better the curve fit.

Figure 4.15. Example of a Berkeley Madonna generated fit. Using data from a batch reaction with a stoichiometry of BnOH/NH3 = 1.

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4.6. Characterisation Techniques

Brunauer-Emmett-Teller (BET) Theory

The Brunauer-Emmett-Teller theory is used to explain the physical adsorption of gas molecules on a solid surface. It is the basis of an analytical technique that measures the specific surface area of solid materials which is commonly used on microporous solid catalysts or supports, including alumina supported catalysts or alumina and clay. In addition, pore volume and pore area distributions in mesopore and macropore range using BJH analysis.

Theoretical Background

The BET theory was developed by Stephen Brunauer, Paul Emmett and Edward Teller.

[143] It describes the phenomenon of multilayer gas adsorption, which is an extension from Langmuir’s theory, among the first to describe a physical adsorption theory for solids. The Langmuir theory describes the relationship between adsorbed gas molecules

(defined as adsorbates) on a solid surface and free gas molecules at the pressure at a fixed temperature:

퐾푃 휃 = (eq. 4.14) 1+(퐾푃)

Where

휃 = fractional cover of the surface

P = gas pressure, in Pa, and

K = equilibrium constant for distribution of adsorbate between the surface and

the gas phase, no units

The Langmuir theory also assumed the following:

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• All surface sites have the same adsorption energy for the adsorbate

• Adsorption of the solvent at one site occurs independently of adsorption at

neighbouring sites

• The activity of adsorbate is directly proportional to its concentration

• Adsorbates form a monolayer, and

• Each active site can be occupied only by one particle

However, there were flaws in the Langmuir theory, which have been addressed by

BET theory with the following assumption:

• Gas molecules will physically adsorb on a solid in layers infinitely

• The different adsorption layers do not interact, and

• The theory can be applied to each layer

By using the Langmuir theory, the BET adsorption isotherm equation can be used to calculate the volume of gas adsorbed at standard temperature and pressure (STP) for an apparent monolayer of gas on the sample surface.

BET adsorption isotherm equation

1 퐶−1 푃 1 (eq. 4.15) 푃0 = × + [푉 ( −1)] 푉푚퐶 푃0 푉푚퐶 푎 푃

Where

• P = partial gas pressure of adsorbate gas in equilibrium with the surface at -

195.8 °C, in Pa

• P0 = saturated pressure of adsorbate gas at the temperature of adsorption, in Pa

• 푉푎 = volume of gas adsorbed at standard temperature and pressure (STP), at 0

°C & 1.1013 x 105 Pa, in cm3

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• 푉푚 = volume of gas adsorbed at standard temperature and pressure (STP) for an

apparent monolayer of gas on the sample surface, in cm3

• 퐶 = dimensionless constant related to the enthalpy of adsorption of the

adsorbate gas on the sample, no unit

푃0 푃 Therefore, by plotting [푉푎 ( − 1)] against , a straight line is obtained at a relative 푃 푃0

퐶−1 pressure of ~ 0.05 - 0.3 Pa. From the resultant linear plot, the gradient ( ) and the 푉푚퐶

1 intercept can be obtained. These values can be used to calculate the volume of gas 푉푚퐶 adsorbed at STP for a monolayer of gas on the sample surface.

Experimental procedures

In this work, the BET adsorption isotherms were determined using nitrogen as a molecular probe gas at -195.8 °C. N2 is relatively inert and easily available, and is ideal for this application, despite liquid N2 occurring at a relatively high temperature. Relative

푃 pressure of ( ) was 0 to 1.0 and the instrument used was a TriStar 3000 (Micromeritics) 푃0 analyser. In a typical measurement, a solid, powder sample (ca. 100 mg) was degassed at 200 °C under flowing nitrogen gas for 4 h using a Flow prep 060 instrument

(Micromeritics), followed by gas evacuation and cooling to -195.8 °C by liquid N2. The sample was then dosed by N2 gas under controlled increments. In between each dose, the pressure was allowed to equilibrate and measured. The quantity of adsorbate (N2) adsorbed on the adsorbent (sample) was recorded and plotted against the equilibrium gas pressure. This data was used to determine the quantity of gas required to saturate the sample surface with a monolayer, and the BET specific surface area (SBET) was calculated using the quantity of gas. The micropore surface area and micropore volume were then calculated by the t-plot method.

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Induced Coupled Plasma Optical Emission Spectroscopy (ICP-OES)

Theoretical Background

Induced coupled plasma optical emission spectroscopy (ICP-OES) is an analytical technique to determine chemical elements in a sample. This is done by vaporising the sampling and exciting atoms and ions with an Argon plasma (~ 7000 °C), causing them to emit electromagnetic radiation at wavelengths characteristic of a particular element.

The radiations are detected with a photomultiplier tube and the separation of spectrum lines is achieved by using a polychromator with a diffraction grating to disperse the light into its component wavelength, and the emission intensity is proportional to the concentration of the element within the sample.

In this work, the ICP-OES elemental analysis was carried out with a PerkinElmer

Optima 2000 DV Inductively Coupled Plasma instrument. Calibrations were carried out using elemental standard solutions (1.0, 2.5 & 10 ppm). The samples were prepared as follows and contained in labelled 13 mL centrifuge tubes for analysis.

Sample preparation

An aqueous solution of 1:3 HNO3 (70 wt%) in HCl (37 wt%), aqua regia, (ca. 20 mL per batch) was prepared with care prior to the digestion of solid and liquid samples.

Solids

Considering a sample of. Ni/Al2O3/SiO2, with a nominal loading of 65 wt% Ni, a small amount of the solid sample (1 mg) was digested in aqua regia (2 mL) at 90 °C for 16 hours with stirring, then the entire digestion was carefully diluted to exactly 10 mL with de-ionised (DI) water. 0.5 mL of the diluted solution was then further diluted to exactly

10 mL with DI water. Assuming all Ni to be dissolved, this sample should contain 6.5 ppm. Complete digestion of the solid samples was ensured so that the metal

120 concentration in solution should be between 1 and 10 ppm. A sample solution of more than 5 mL is required for an accurate reading.

Liquids

Ni catalyst extrudites (Catalyst 2) with a nominal loading of 36 wt% Ni (293.5 mg, 1.8 mmol Ni) were used as a catalyst in 30 ml organic solution, o-xylene, which was heated for 72 hours at 160 °C. To test for the Ni content in the organic solution, the procedures detailed below were followed:

5 ml of solution was concentrated under vacuo at 70 °C for 15 minutes. The residue was digested with aqua regia (0.4 ml) at 90 °C for 16 hours with stirring. The entire digestion was carefully diluted to exactly 10 mL with DI water. 0.2 mL of the diluted stock was then further diluted to exactly 10 mL with DI water.

Assuming all the Ni to be dissolved in the organic phase, the sample should contain 35 ppm. In reality, the Ni concentration will be much lower, as it is not expected that all Ni content will be leached from the catalyst.

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Transmission Electron Microscopy (TEM-EDX)

Theoretical background

Transmission electron microscopy (TEM) is a powerful analytical technique for structural characterisation. An important application is the atomic-resolution real-space imaging of nanoparticles, which allows determination of particle shapes and lattice structures. In practice, TEM examines structure by passing electrons through the specimen, and the image is formed as a shadow on a phosphorescent screen. In order for the electrons to pass through, a very thin layer (< 100 nm) of the sample is necessary to be dispersed onto copper grids. TEM, in tandem with energy-dispersive x-ray spectrometer (EDX or EDS), is a powerful tool that provides chemical information at a spatial resolution of 1 nm or better.

Energy-dispersive x-ray spectrometer (EDX or EDS) is a spectroscopic technique that detects x-rays emitted from samples during bombardment by an electron beam to characterise the elemental composition of the samples. The X-ray detector measures the relative abundance of emitted x-rays against their energies, and the spectrum of X-ray energy versus counts is evaluated to determine the elemental composition of the samples.

Experimental Procedure

A JEOL 2000-FX 200KV microscope coupled with an energy-dispersive X-ray spectrometer (EDX) was used. A small amount of sample (ca. 1 mg) was dispersed in ethanol prior to depositing onto holey carbon films on 300 mesh copper grids (Agar

Scientific). The grids were allowed to dry under room temperature and pressure for 20 minutes until all ethanol evaporated. The grids were then placed in sample holders for transportation.

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Powder X-Ray Diffraction (XRD)

Powder X-Ray diffraction spectroscopy is a rapid analytical technique primarily used for phase identification of crystalline material, and has been used for catalyst characterisation. This technique is fast and non-destructive and investigates the unit cell dimensions of the sample and the catalyst can still be used for further reactions [144].

Theoretical background

In 1912, Max von Laue discovered that crystalline substances act as a 3-D diffraction grating for X-ray wavelengths similar to the spacing of planes in a crystal lattice (ca.

1Å), for which he won the Nobel Prize in Physics in 1914 [145]. More specifically, when the crystalline material is exposed to monochromatic x-rays, the incident beam is diffracted at specific angles to the lattice planes and can be detected as a characteristic pattern, known as the Bragg diffraction peaks, using formulae developed by William

Henry Bragg and William Lawrence Bragg, who later shared the Nobel Prize in Physics in 1915. [146] X-ray diffraction is the most powerful for crystalline materials, as the long-range order will always show Bragg diffraction peaks, and structural disorder or defects may lead to scattering and disappearance of the peaks [144].

Bragg’s law governs the conditions for the diffraction:

푛휆 = 2푑 sin 휃 (eq. 4.18)

Where:

n = an integer

λ = wavelength of the incident X-ray radiation, in nm

d = distance between the atomic layers in a crystal, or the d-spacing, and

θ = angle of incidence between the X-ray beam and the diffraction planes, in °

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In 1918, Paul Scherrer determined the inverse relationship between the width of an X- ray diffraction peak and the sample crystalline size. [147] The crystalline sizes can be estimated with the Scherrer equation,

Kλ D = (eq. 4.19) βcosθ

Where:

D = mean grain size

K = dimensionless shape factor = 0.9

λ = wavelength of incident X-ray radiation (with Cu anode, λ = 0.154 nm)

1/2 β = 0.5 H (π / loge2) (Assuming Gaussian peak)

H = Full-Width Half Maximum (FWHM)

θ = Bragg angle, i.e., angle at maximum peak (in radians)

Experimental Procedure

XRD was used for qualitative identification of crystalline phases in samples, as well as the determination of particle sizes. A finely ground, homogenised average bulk composition sample must be used to achieve repeatability. The equipment used was an

X’pert PRO Diffractometer (PANalytical) using CuKα radiation (α = 1.5404 Å) with an

X’Celerator detector with Ni filter and Soller slit system. Measurements were carried out at RTP (22 °C, 1 bar). Data analysis was carried out with PANalytical X’pert High

Score Plus software. The instrument broadening of standard Si (111) at 2θ = 28.44 and the CuKα1 radiation used was 0.15.

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Temperature programmed desorption (TPD)

Theoretical background

Temperature programmed desorption (TPD) has been widely used for characterising the acid and basic sites on catalyst surfaces. Determining the strengths and quantity of the sites is important for understanding the performance of catalysts. This is done by investigating the temperature-dependence of specific adsorption or desorption processes of samples in a well-defined gas atmosphere. NH3 and CO2 are commonly used as probing gases for acid and basic sites respectively. Other probing agents include pyridine and t-butyl alcohol for acid sites and pyrrole for basic sites. [148]

The detection of probe gases was done by mass spectrometry. Mass spectrometry is a versatile analytical technique widely used for both quantitative and qualitative analysis of species in a sample by measuring the mass-to-charge ratio and abundance of gas- phase ions.

Experimental setup

The equipment components are shown in Table 4.6. The flow diagram of how the components are connected is shown in Figure 4.16.

Table 4.6. Components of the Temperature Programmed Desorption equipment.

Component Details Purpose

Gas cylinders NH3, CO2, He Gas supply

Mass flow controllers Bronkhorst EL-Flow Gas flow control

Furnace HY-800 Shinch Sample heater

Temperature controller West 2400 DIN Temperature control

Mass Spectrometer Fisons SXP HR-QMS Gas composition measurements

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Figure 4.16. Temperature programmed desorption equipment setup flow diagram. The components shown in the images are listed in Table 4.6.

Experimental Procedure

Sample preparation

0.20 g of dry sample (e.g. Ni/Al2O3/SiO2) was placed in the middle of an O.D. ¼’’ quartz tube, held in place by quartz wool (ca. 0.1 g). The sample was subjected to a constant flow of He at 30 mL/min to prevent oxidation/contamination and was degassed at 100

°C for one hour to remove water vapour while avoiding pore damage. This was followed by a heat ramp up to 500 °C at 10 °C/min to remove any strongly adsorbed species on the catalyst. The catalyst was then allowed to cool to 100 °C under He.

Adsorption

3 mL/min of basic or acidic probe gas was added to the gas flow, (ammonia or carbon dioxide respectively), whch was diluted by the constant He flow of 30 mL/min. This amounted to a dilution of 10% v/v. The temperature was then set at 100 °C for 1 hour.

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The probe gas was then switched off and the sample remained at 100 °C for an additional hour to remove any physically adsorbed probe gas.

Desorption

Starting from 100 °C, the sample was heated to 750 °C at a rate of 10 °C/min, and held at 750 °C for 30 minutes. The gas composition was analysed by a mass spectrometer calibrated to the gases prior to the experiment.

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Thermogravimetric analysis – Mass spectrometry (TGA – MS)

Theoretical Background

Thermogravimetric analysis (TGA) is the measurement of specific mass difference when a sample is subjected to controlled temperature, pressure and atmosphere. The evolved gas can then be analysed by mass spectrometry (MS). This change in the specific mass is benchmarked with respect to an inert reference sample in an identical thermal regime. This allows a continuous determination of sample weight as a function of time/temperature in various atmospheres. The resultant fumes are then carried into a mass spectrometer, where the masses of the gaseous fragments will be detected over time/temperature. Quantification of gases (H2O & CO2, using the decomposition of

CaCO3) can be achieved by calibration with known amounts of substance heated over time.

Experimental procedure

The equipment used for TGA was the TA-Q500 coupled with a Discovery mass spectrometer. Ca. 20 mg of samples were loaded onto the platinum trays for each run.

A flow rate of 200 mL/min 10% air/argon (v/v%) mixture was used at a heating rate of

20 °C/min from room temperature up to 500 °C. The change in sample mass and intensities of fragments are recorded over time and presented as a function of temperature.

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Gas chromatography – Flame Ionisation Detection (GC-FID)

Theoretical background

Gas chromatography (GC) is a common analytical technique in which samples (usually a mixture of components) are volatilised without decomposition, then interacts with the surface of the column which is coated in a stationary phase. [149] The components elute at different times, known as the retention times of the components. The components are then detected with a detector as they exit the column, which generates a signal that can be used to create calibration curves.

In this work, a flame ionisation detector (FID) was used, which pyrolyses the components, generating an increase in current between electrodes placed next to the flames. The increase in current is then translated which generates peaks on the chromatograph.

Experimental procedure

Alcohol conversions and amine selectivities were determined by comparison with known standards using an HP 5890 gas chromatography system fitted with a flame ionisation detector, using calibration plots for reactant and products, against 4-tert- butylphenol as an internal standard. A 0.1 M 4-tert-butylphenol (4 – tBuPhenol) solution in methanol prepared beforehand and was added to the sample as a dilutant before injection to the GC. 4-tert butylphenol is used as an internal standard for its high availability, as well as its solubility in o-xylene and its inertness with reactants and potential products.

The products were further confirmed by a 1380 GC system fitted with mass spectrometry (MS). Both GC systems were fitted with HP-5 columns (25 m x 0.53 mm,

5.00 µm), using nitrogen and helium as carrier gases respectively. The injector temperature was 250 °C and detector 280 °C. The heating ramp began at 50 °C,

129 isothermal for 2 minutes, followed by a ramp up at 50 °C/min to 300 °C, and stayed at

300 °C for 5 minutes. The retention times of common species detected by GC-FID is shown below in Table 4.6.

Table 4.7. Retention times of common species detected by GC-FID

Component Retention time (min) Calibration factor

Methanol (dilutant) 1.2 /

Toluene 1.6 1.7

o-xylene (solvent) 1.6 /

Benzaldehyde 3.16-3.19 6.7

BnNH2 3.33-3.37 7.6

BnOH 3.44-3.48 7.7

4-tert-butylphenol 4.71 1

Bn2NH 5.17-5.20 13.7

PhCHNBn 5.25-5.27 13.8

Bn3N 6.96 20

Solution preparation:

100 µL of the reaction solution was transferred to an HPLC vial (1 mL), followed by

500 µL of a dilute 4-tert-butylphenol solution in methanol (0.10 M). The vial was then sealed and shaken thoroughly to ensure complete mixing before submitting to GC analysis.

Calibration curves

The calibration curves were carried out at known concentrations of standard components. The ratio between the areas of the component and 4-tert-butylphenol

(standard) was plotted against the concentrations. The gradient of this curve was then used as the calibration factor. One example of Bn2NH is shown in Figure 4.16.

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Figure 4.17. Calibration graph for Bn2NH. Concentrations of Bn2NH solutions were analysed by GC-FID (x-axis) against (area of the component)/(area of the standard) (y- axis)

4.7. Conclusion

In this chapter, all materials used in the project and their suppliers were illustrated.

Experimental methods, including batch and flow reactions and catalyst preparations, computational and analytical methods used in this project are explained, which will be referred to in later chapters and sections.

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Chapter 5: Batch Reactions

5.1. Introduction

In this chapter, the N-alkylation of ammonia (NH3) and benzyl alcohol (BnOH) using hydrogen borrowing cycle are performed in batch reactions. The following aspects will be investigated:

• thermodynamics of the reaction will first be calculated,

• preliminary batch reactions with catalyst screening,

• kinetic modelling using experimental data,

• further reactions with reaction intermediates and refining the kinetic model,

• the effect of water,

• temperature studies,

• exploring the scope of alcohol substrates,

• comparison of this reaction with literature.

BnOH is commonly used as a model compound for alcohol to amine conversions due to its availability. Therefore, using BnOH allows easier comparison of the current reaction with other literature, as many research groups had selected BnOH as the model alcohol in their work. The hydrogen borrowing cycle of this reaction is shown in

Scheme 5.1.

Scheme 5.1. Scheme for hydrogen borrowing cycle of NH3 alkylation with BnOH, where R = Ph.

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The hydrogen borrowing cycle, also known as the hydrogen auto-transfer reaction, involves 3 steps:

1) Benzyl alcohol (BnOH) undergoes oxidation to benzaldehyde (PhCHO),

2) benzaldehyde then goes through a condensation reaction with ammonia

(NH3) to give benzamidine (PhCHNH), and

3) benzamidine subsequently undergoes reduction by the initially generated hydrogen to yield the benzylamine (BnNH2).

The hydrogen is effectively “borrowed” from the starting alcohol and “auto-transferred” to the imine to produce the final amine, and hence the name of the mechanism.

The reaction to produce BnNH2 can be simplified, as shown in Scheme 5.2.

Scheme 5.2. Primary amine formation from BnOH+NH3.

However, there is a major challenge in this reaction, namely the possibility of further reaction between BnNH2 and BnOH. This is because BnNH2 is a primary amine, and is more nucleophilic than NH3. The product of this reaction is dibenzyl amine (Bn2NH), which may further react with BnOH to form the tertiary amine, tribenzyl amine (Bn3N).

These reactions are shown in Scheme 5.3.

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Scheme 5.3. Expected Bn2NH and Bn3N formation from BnOH+NH3 reaction.

These reactions are combined and simplified and shown in Scheme 5.4.

Scheme 5.4. Possible reaction products with BnOH and NH3 are BnNH2, Bn2NH and

Bn3N.

As previously discussed in Chapter 2.7, noble metals, such as Au and Ru, achieved high selectivities to the higher amines [121]; whereas base metals such as Ni and Cu are more selective towards the primary amine [125]. Moreover, the stoichiometry of reactants has a large impact on the selectivities of product amines: the higher the NH3 to alcohol ratio, the higher the selectivity to the primary amine [121, 122]. Therefore, conditions must be carefully controlled to achieve the desired yields.

This chapter focuses on the N-alkylation of NH3 with BnOH carried out under batch conditions, where

• The thermodynamics of the system was first predicted under different

temperature, pressure and initial concentrations.

• Preliminary reactions were carried out using different commercial catalysts to

compare performances.

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• Experiments with different water concentrations and NH3/BnOH ratios were

carried out, and a condensation method was developed to eliminate the use of

water in all reactions.

• Experimental data were simulated and analysed using reaction kinetics to

improve on the current results and to better understand how the reaction

parameters affected the system

• Batch reactions between NH3 and several primary and secondary alcohols were

demonstrated, broadening the substrate scope.

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5.2. Thermodynamics

Thermodynamics considers the state of the system by comparing the energy differences between the products and reactants produced during chemical reactions. The Gibbs free energy (ΔG) is the thermodynamic potential of the system that can be used to calculate the maximum reversible work that may be performed by the thermodynamic system under a constant temperature and pressure. It can be calculated by:

훥퐺 = 훥퐻 − 푇훥푆 (eq. 5.1)

Where

• ΔG is the Gibbs free energy of the system, measured in kJ;

• ΔH is the enthalpy difference between products and reactants, measured in kJ;

• T is the absolute temperature, measured in K; and

• ΔS is the entropy change between products and reactants, measured in kJ K-1.

When ΔG is negative, the reaction is thermodynamically favoured and will occur spontaneously; when ΔG is positive, the reaction is non-spontaneous and additional energy is required for the forward reaction to proceed; and when ΔG is zero, the process is in equilibrium.

However, thermodynamics gives information of the possibility of a reaction happening as well as the equilibrium position of products at the end of the reaction, based on the reaction conditions, but gives no insight into how quickly this equilibrium is reached.

The rate at which a reaction reaches equilibrium is dictated by reaction kinetics, which will be discussed later in this chapter.

Once the standard free energy change, ΔG, is known for a chemical process at temperature T, the equilibrium constant, K, for the overall process at that temperature can be calculated using the following equation:

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훥퐺 = −푅푇푙푛퐾 (eq. 5.2)

Where R is the universal gas constant (8.314 J K-1 mol-1) and the temperature is measured in Kelvin (K). This equation can be rearranged to give:

퐾 = 푒훥퐺/(−푅푇) (eq. 5.3)

In the model reaction (BnOH + NH3), there are three main products that can potentially be produced. The overall reaction scheme is given below as Scheme 5.5. This reaction will be known as reaction A.

Scheme 5.5. Reaction A: overall reaction scheme for BnOH+NH3.

Each reaction step can be broken down into the three following steps, as illustrated in

Schemes 5.6, 5.7 and 5.8:

Scheme 5.6. Reaction 1: The first step of reaction A, where BnOH reacts with NH3.

[퐵푛푂퐻] + [푁퐻3] ↔ [퐵푛푁퐻2] + [퐻2푂] (Reaction 1)

[퐵푛푁퐻2][퐻2푂] 퐾1 = (Reaction 1) [퐵푛푂퐻][푁퐻3]

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Scheme 5.7. Reaction 2: The second step of reaction A, where BnOH reacts with

BnNH2.

[퐵푛푁퐻2] + [퐵푛푂퐻] ↔ [퐵푛2푁퐻] + [퐻2푂] (Reaction 2)

[퐵푛2푁퐻][퐻2푂] 퐾2 = (Reaction 2) [퐵푛푁퐻2][퐵푛푂퐻]

Scheme 5.8. Reaction 3: The third step of reaction A, where BnOH reacts with

Bn2NH.

[퐵푛2푁퐻] + [퐵푛푂퐻] ↔ [퐵푛3푁] + [퐻2푂] (Reaction 3)

[퐵푛3푁]+[퐻2푂] 퐾3 = (Reaction 3) [퐵푛2푁퐻]+[퐵푛푂퐻]

Values of the equilibrium constants K1, K2 and K3 were determined by the overall ΔG of reactions using ASPEN-plus. The modelling software can calculate the equilibrium position of the overall process with a simple equilibrium reactor. By defining different conditions, such as pressure, temperature and reactant stoichiometry, the reaction outcome can be predicted.

The equation of state property method used in ASPEN-plus was the Electrolyte Non-

Random Two-Liquid model in tandem with the Redlich-Kwong equation of state

(ENRTL-RK). This method was used instead of the IDEAL EOS because although both are good predictive models for systems at any concentration under high pressures,

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ENRTL-RK allows the prediction for mixtures of polar and non-polar compounds. This is highly relevant because there are polar (water and ammonia) and non-polar (o-xylene) components in this system [150]. The simple RK method was also considered, but

ENTRL-RK was better suited due to its considerations in non-ideal conditions with high pressure and temperature.

The reaction conditions were first set at a pressure of 10 bar and temperature of 160 °C, then at pressure 60 bar and temperatures of 160, 120 and 80 °C. The different sets of temperature and pressure are selected to proof the following concepts. Temperature may affect the equilibrium positions of these reactions, as higher the temperature higher is the entropy. On the other hand, pressure should not affect the equilibrium position because these reactions happen in the liquid phase and any pressure changes should not affect the overall entropy of the system. These conditions were tested with different

BnOH:NH3:H2O ratios as follows: 1:1:1, 1:1:10, 1:1:100 & 1:1:0 to compare the effect of excess H2O, a shift in equilibrium is expected to happen with more water. For anhydrous reactions, 6:1:0 is an excess of BnOH, where a high production of Bn3N is expected, and 1:3:0 is an excess of NH3, where a high production of BnNH2 is expected.

Table 5.1 below shows the results from the thermodynamics prediction. The units for input and output streams were kmol/hr and are shown in grey and white boxes respectively. The conversion based on converted BnOH and selectivity to the products are shown in italics.

The conversions are calculated with respect to the BnOH consumed in the reaction, following the equation below:

[퐵푛푂퐻] −[퐵푛푂퐻] 퐶표푛푣푒푟푠𝑖표푛 = 푖푛푝푢푡 표푢푡푝푢푡 × 100% (eq. 5.4) [퐵푛푂퐻]푖푛푝푢푡

The selectivity of a product is defined as the ratio between the product and all the products, described by the equation below:

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[퐵푛푁퐻2] 푆푒푙푒푐푡𝑖푣𝑖푡푦 표푓 퐵푛푁퐻2 = × 100% (eq. 5.5) [퐵푛푁퐻2]+[퐵푛2푁퐻]×2+[퐵푛3푁]×3

[퐵푛2푁퐻]×2 푆푒푙푒푐푡𝑖푣𝑖푡푦 표푓 퐵푛2푁퐻 = × 100% (eq. 5.6) [퐵푛푁퐻2]+[퐵푛2푁퐻]×2+[퐵푛3푁]×3

[퐵푛3푁]×3 푆푒푙푒푐푡𝑖푣𝑖푡푦 표푓 퐵푛3푁 = × 100% (eq. 5.7) [퐵푛푁퐻2]+[퐵푛2푁퐻]×2+[퐵푛3푁]×3

In the calculations, the concentrations of Bn2NH and Bn3N are multiplied by 2 and 3

respectively because [Bn] is considered as one unit.

Table 5.1. Thermodynamics data predicted by ASPEN-plus with ENTRL-RK method.

Conditions BnOH NH3 H2O RNH2 R2NH R3N Reaction Reaction Temperature Pressure Input stream (kmol/hr) Sel.(%) Conv. Of Entry (°C) (Bar) Output stream (kmol/hr) BnOH (%)

1 1 1 80 3 17 100 1 0 0.13 2 0.80 0.02 0.06

1 1 10 80 3 17 100 2 0 0.13 11 0.80 0.02 0.06

1 1 100 81 3 16 99 3 160 10 0.01 0.13 101 0.80 0.02 0.05

6 1 0 0 0 100 50 4 3 0 3 0 0 1

1 3 0 100 0 0 100 5 1 2 1 1 0 0

1 1 1 80 3 17 100 6 160 60 0 0.13 2 0.80 0.02 0.06

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1 1 10 80 3 16 100 7 0 0.13 11 0.80 0.02 0.06

1 1 0 82 3 15 100 8 0 0.12 1 0.82 0.02 0.05

6 1 0 0 0 100 50 9 3 0 3 0 0 1

1 3 0 100 0 0 100 10 1 2 1 1 0 0

1 1 1 83 2 14 100 11 0 0.13 11 0.80 0.02 0.06

1 1 10 83 2 14 100 12 0 0.13 11 0.80 0.02 0.06

1 1 0 86 2 12 100 13 120 60 0 0.09 1 0.87 0.01 0.04

6 1 0 0 0 100 50 14 3 0 3 0 0 1

1 3 0 100 0 0 100 15 1 2 1 1 0 0

1 1 1 84 2 14 100 16 0 0.1 2 0.84 0.01 0.05

80 60 1 1 10 84 2 14 100 17 0 0.13 11 0.80 0.02 0.06

18 1 1 0 86 2 12 100

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0 0.09 1 0.87 0.01 0.04

6 1 0 0 0 100 50 19 3 0 3 0 0 1

1 3 0 100 0 0 100 20 1 2 1 1 0 0

It was observed that different pressures did not affect the reaction equilibrium. The

results were identical when the pressure was increased from 10 to 60 bar (Table 5.1,

Entries 1-5, Entries 6-10). This was expected because the reaction happened in liquid

phase, where the pressure changes would have no effect to the entropy of the system.

However, decreasing the temperature from 160 °C to 80 °C and 120 °C had a small

effect on the equilibrium, namely a decrease of 0.01 kmol/hr in Bn3N and an increase

of 0.03 kmol/hr in BnOH at lower temperatures (Table 5.1. Entries 8, 13, 18).

A conversion of 100% can be observed in many entries, which is expected because the

thermodynamic calculations consider the outcome of completed reactions. However, at

a high excess of water (Table 5.1, Entry 3), where the BnOH/NH3/H2O ratio was

1:1:100, the conversion was 99%. The incomplete conversion suggests that water

inhibits the overall reaction, as evident in the hydrogen borrowing cycle, where the

condensation step is reversible, as shown in Scheme 5.9.

Scheme 5.9. Hydrogen borrowing cycle with emphasis on the condensation step.

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Product Selectivities

The reaction pressure had very little effect on product selectivity, which is expected because the reaction happened in liquid phase, and pressure changes had no effect to the entropy changes to the system.

There are three possible products in this reaction, BnNH2, Bn2NH and Bn3N. For reactions where BnOH:NH3 ratio = 1:1 (Table 5.1. Entries 1 – 3, 6 – 8, 11 – 13, 16 –

18), the selectivity to BnNH2 was the highest at 80 – 87%, with 12 – 17% Bn3N and 2

– 3% Bn2NH across 80 – 160 °C. Therefore, BnNH2 is expected to be the major product when equimolar of BnOH and NH3 are used.

BnOH > NH3

When an excess of BnOH was used at BnOH:NH3 ratio = 6:1 (Table 5.1. Entries 4, 9,

14, 19), 100% selectivity to Bn3N was observed, which is expected because the system would proceed to form the most stable product, which is Bn3N.

BnOH < NH3

When an excess of NH3 was used with BnOH:NH3 = 1:3 (Table 5.1. Entries 5, 10, 15,

20), 100% selectivity to BnNH2 was observed, which is also expected because the excess NH3 would limit the product formation to only BnNH2.

In summary, thermodynamics outcomes calculated by ASPEN-plus considered the outcome at reaction completion. The pressures and temperatures had very little effect on the reaction outcome, and in most cases, BnOH was completely consumed. The stoichiometry of the reactants played a major role in product selectivity, which can be controlled by increasing either BnOH or NH3. When water was used in excess, it had a minute effect on the overall outcome.

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5.3. Preliminary Batch Experiments with Commercial Catalysts

In recent research, heterogeneous catalysts were developed for the synthesis of amines using alcohols and NH3, including Ru, Ni, Cu and Fe. Such examples were discussed in

Chapter 2 in full detail. However, only a few examples demonstrated high selectivities towards primary amines, which included Ni/Al2O3 at 90 – 100% [122]; whereas the Cu and Ru catalysts had higher yields to secondary and tertiary amines respectively at 80% and 87% respectively [120][121]. However, these reactions were carried out under different conditions, with temperatures ranging from 131 to 160 °C, different reactant stoichiometry at NH3/BnOH = 0.2 to 2.2 and different reaction times of 6 to 72 hours.

Scheme 5.10. Examples of N-alkylation of NH3 with BnOH that resulted in different selectivities.

The reaction conditions (temperature, reactant stoichiometry and pressure) across these experiments were not consistent and were thus difficult to compare. Considering this, the following preliminary reactions were designed to have consistent stoichiometries at

NH3/BnOH = 1, 160 °C, initial 4 bar N2, 30 mL of 0.1 M BnOH in a pressurised batch reactor at 1200 rpm, at 4 hours.

The batch reactor had a headspace of 10 mL, as well as a sample dispenser which allowed sampling over time (Chapter 4, Section 4.3.1, Figure 4.1). The results were analysed by GC-FID using an HP5890 fitted with an HP5 column. Each sample (100

144

µL) were diluted using 0.01 M 4-tert-butylphenol solution in methanol (500 µL). If necessary, the samples were analysed by a 1380 GC-MS system fitted with an HP5 column.

A number of commercially available noble metal and nickel catalysts were selected and their activities tested to determine which one(s) would achieve high selectivity to primary amine under the same reactant stoichiometry. These included 65 wt%

Ni/Al2O3/SiO2, 16 wt% Ni/Al2O3/CaO, 16 wt% NiMo/Al2O3, 1 wt% Au/TiO2, 1 wt%

Ru/Al2O3 and 1 wt% Pt/Al2O3.

All of these catalysts were selected for their potential activity in the alkylation of amines via the hydrogen borrowing cycle, as well as other attractive features for this work.

The Ni catalysts were selected for their robustness and long lifetime as they have applications in hydrogenation and hydrotreatment processes [151]. The robustness is an attractive feature because when applied to a continuous reactor the catalyst’s lifetime would directly impact the running cost of the overall process [152]. The noble metal catalysts (Au, Ru & Pt) are highly efficient in hydrogenation as well as amine alkylation reactions via the hydrogen borrowing cycle, so it would be interesting to see their activity in NH3 alkylation [153]. However, Au and Pt had not been shown to catalyse

NH3 alkylation with alcohols. These catalyst supports were chosen for their surface basicity and acidity for interactions between the alcohols and amines respectively.

Samples were analysed by GC-FID using calibration curves of reactants and potential products.

Initial batch reactions were conducted with 30 mL batch reactors. All catalysts were used as purchased with no pre-treatment prior to the experiment. 30 mL of 0.1 M BnOH solution in o-xylene was used, as o-xylene is an inert solvent. Toluene was also considered as a potential solvent, but it has a lower boiling point of 110.6 °C and could be a potential side product in the disproportionation reaction of BnOH.

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A low concentration of 0.1 M was used to decrease the likelihood of disproportionation, where BnOH reacts with itself to form toluene and benzaldehyde. Fang et al. [154] reported that high concentrations of BnOH are prone to disproportionation reactions, especially with noble metal catalysts. Aqueous NH3 was used in these preliminary experiments due to the ease of NH3 quantification, as well as the safety in handling and transportation of the solution. Imines were not detected in these experiments. 4 bar of

N2 was purged and charged in the reactor prior to the reaction to ensure an inert gaseous environment. A temperature of 160 °C was used as it is slightly above the boiling point of o-xylene (b.p. = 144 °C). Catalyst loadings of Ni 5 mol% and M 1 mol% were used, where M represents noble metals. The results of the preliminary batch reactions are shown in Table 5.2.

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Table 5.2. Preliminary batch results with commercial catalysts a

Conv. Sel. (%)

Entry Catalyst (%) BnNH2 Bn2NH Bn3N PhCHO

b 1 65 wt% Ni/Al2O3/SiO2 44 85 15 0 0

b 2 16 wt% Ni/Al2O3/CaO 20 96 0 0 3

b 3 16 wt% NiMo/Al2O3 10 0 0 0 100

c 4 1 wt% Au/TiO2 48 0 60 37 0

c 5 1 wt% Ru/Al2O3 29 0 12 63 20

c 6 1 wt% Pt/Al2O3 16 12 32 36 16

7 Thermodynamic prediction 100 84 1 15 0

a Conditions: 0.10 M BnOH in o-xylene, 30 mL, 160 °C, 1200 rpm, initial 4 bar N2, 4 hours,

b c BnOH/NH3/H2O = 1/1/5. Samples were analysed by GC-FID. ROH/Ni = 20, ROH/M = 100

At first sight, both Ni catalysts (65 wt% and 16 wt%) were more selective towards primary amine at 85% and 96% respectively (Table 5.2, entries 1 & 2). This high selectivity to primary amine was expected as demonstrated by previous studies of other groups [122]. 65 wt% Ni/Al2O3/SiO2 had a 44% conversion, which is the highest among the nickel catalysts. An 85% selectivity towards BnNH2 was observed, and Bn2NH at

15%. No Bn3N or other side products were observed. These results were similar to those reported by Shimizu et al. with a 10%wt Ni/Al2O3 catalyst at 60% conversion and 90% selectivity [155]. However, the reaction by Shimizu et al. [155] had a longer reaction time of 24 hours, compared to 4 hours here, and a higher NH3/BnOH of 2.2 compared to 1 here. Moreover, their reaction used gaseous NH3 while aqueous NH3 was used in these preliminary batch reactions. However, in this reaction, a higher Ni loading of 5% was used compared to 1% by Shimizu et al. [155], which may explain the faster reaction rate.

16 wt% Ni/Al2O3/CaO catalyst (Table 5.2, entry 2) gave a lower conversion at 20% but the highest selectivity towards primary amine at 96%. However, the commercial Ni/Mo

147 catalyst, 16 wt% NiMo/Al2O3 (Table 5.2, entry 3) saw little conversion at 10% with only benzaldehyde as the product. The low conversion could be due to NiO present in the catalyst, which failed to reduce the imines to amines in the hydrogen borrowing cycle. The imine then reverted back to aldehyde, as aldehyde and NH3 is a more thermodynamically stable combination compared to imine and water (Scheme 5.7).

Scheme 5.11. Proposed scheme of formation of aldehyde due to an incomplete hydrogen borrowing cycle.

The presence of NiO is confirmed by XRD. This is directly compared with 65 wt%

Ni/Al2O3/SiO2 catalyst, as shown in Figure 5.1. The peak at 43.29 ° (2θ) corresponds to the NiO [2 0 0] plane, whereas the peak at 44.50 ° (2θ) corresponds to the Ni [1 1 1] plane. This matches with the values reported by Richardson et al. [137] of 43.30 ° and

44.50 ° respectively.

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Ni [1 1 1]

65 wt% Ni/Al2O3/SiO2

NiO [2 0 0]

16 wt% NiMo/Al2O3

Figure 5.1. XRD patterns for the commercial catalysts, 65 wt% Ni/Al2O3/SiO2 and 16 wt% NiMo/Al2O3.

On the other hand, the noble metal catalysts were more selective towards secondary and tertiary amines, as suggested in the literature [121]. Au/TiO2 (Table 5.2, entry 4) had a higher conversion of 48% compared to Ni and other noble metals, with selectivities of

60% and 37% to secondary and tertiary amines respectively. This agrees with the literature, as the Au/TiO2 catalyst had previously been reported to be highly active in the N-alkylation of aniline with alcohols under flow conditions to form secondary amines [25]. The reaction scheme is shown below in Scheme 5.12.

Scheme 5.12. Reaction scheme of BnOH and aniline forming secondary amine.

Ru/Al2O3 (Table 5.2, entry 5) gave a conversion of 29% with selectivities towards the secondary and tertiary amines at 12% and 63% respectively. Pt/Al2O3 (Table 5.2, entry

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6) gave a low conversion of 16% with low selectivity to primary amine at 12%, and moderate selectivities to Bn2NH and Bn3N at 32% and 36% respectively.

Benzaldehyde was detected for reactions with Ru and Pt catalysts (Table 5.2, entries 5

& 6). Benzaldehyde is known to be one of the two products in the disproportionation reaction of BnOH (Scheme 5.13). Toluene was not detected as it overlaps with the solvent peak in GC analysis. This agrees with other works that involved high concentrations of benzyl alcohol and noble metal catalysis. For instance, Morad et al.

[156] reported disproportionation reaction competed with the aerobic oxidation of solventless BnOH in a flow reactor using the 1% AuPd/TiO2 catalyst, where the reaction afforded 21% selectivity towards toluene.

Scheme 5.13. Disproportionation reaction of benzyl alcohol to toluene and benzaldehyde.

When compared with the theoretical results calculated by thermodynamics in the previous chapter, there is clearly a difference between the experimental and theoretical results. The theoretical results indicated a 100% conversion of BnOH to the product amines, as they assume the system reacted to completion; in the experimental results, the reactions were only carried out for 4 hours and stopped before completion, thus only achieving 10 – 44% conversions. The theoretical results indicated a high selectivity of

84% to BnNH2, followed by 15% of Bn3N and only 1% of Bn2NH. 65 wt%

Ni/Al2O3/SiO2 achieved a high selectivity to BnNH2 at 85%, which match the theoretical results; however, the catalyst also achieved a selectivity of 15% to Bn2NH, which does not match the theoretical results. On the other hand, the Au and Ru catalysts demonstrated high selectivities to Bn2NH and Bn3N at 60% and 63% respectively. As

150 none of the experimental results matched the theoretical thermodynamic results, it suggests that the reactions were kinetically controlled by the catalysts. Therefore the reaction kinetics must be investigated.

In summary, the commercial catalysts, 65 wt%Ni/Al2O3/SiO2 and 1wt%Au/TiO2 showed moderate BnOH conversions at 44 – 48% with very different product selectivities. Ni was more selective to BnNH2 at 96% while Au was selective to Bn2NH at 60%. Further investigations will be carried out for these experiments in terms of their reaction kinetics.

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5.4. Kinetic Studies

Reaction kinetics is the study of how different conditions influence the rate of a chemical reaction, which allows insight to the mechanism and/or transition states.

However, kinetics does not give information about the reaction outcome, which is determined by thermodynamics. By studying the reaction kinetics in this project, changes can be made to experimental conditions to achieve higher conversions at a shorter time, as well as higher selectivities to the primary amine by reducing the formation of side products.

A kinetic model is a series of differential equations that describe changes in substrate and product concentrations over time. These differential equations can be solved, and concentrations and reaction rates can be calculated, thus simulating the reaction profiles.

The construction of such mathematical models to describe chemical reactions could be aided by Berkeley Madonna, an ordinary differential equation solver with curve-fitting tools. By understanding the reaction kinetics of the current reaction, one could control the outcome of reactions to increase selectivity to the primary amine.

The fourth-order Runge-Kutta method (RK4) was used for the calculations in Berkeley

Madonna. RK4 belongs to a family of mathematical methods that numerically integrate ordinary differential equations by using four approximations to the slope. This method is preferable over the Euler method because RK4 is more accurate, as it considers four derivatives instead of only one. RK4 is a relatively quick method which saves computational time. The full mathematical method is illustrated in the appendix (S3).

The wellness of a curve fit is determined by the root mean square (RMS) deviation between the data and the best run. This deviation is calculated using the root mean square of the differences between individual data points in the dataset and the corresponding points in the run. The lower the RMS value in this software, the better the curve fit.

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It is important to note that these kinetic rates are not absolute rate constants for these reactions, as the concentration of the catalyst has been included in the calculations.

5.4.1. Kinetic Models with Ni and Au Catalysed Reactions

In this kinetic model, the overall reactions were simplified to three reaction steps, as shown in Scheme 5.14.

Scheme 5.14. A three-reaction kinetic model in the NH3 N-alkylation reaction with

BnOH.

The reaction rate constants are defined as follows:

• k1f = the forward kinetic rate constant for the 1st reaction,

• k1r = the reverse kinetic rate constant for the 1st reaction;

• k2f = the forward kinetic rate constant for the 2nd reaction,

• k2r = the reverse kinetic rate constant for the 2nd reaction;

• k3f = the forward kinetic rate constant for the 3rd reaction, and

• k3r = the reverse kinetic rate constant for the 3rd reaction

Each reaction was treated as a bimolecular system. The differential equations developed for this model are shown in the appendix. (S4)

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The reaction rate constants, k1f, k1r, k2f, k2r, k3f and k3r were then fitted using experimental data from the two batch reactions involving commercial 65 wt%

Ni/Al2O3/SiO2 and 1wt% Au/TiO2 catalysts. These two reactions were chosen for their high conversions as well as high selectivities to primary and secondary amine respectively. k, the reaction rate constant, quantifies the rate of a chemical reaction. For instance, for a reaction where A+B = P, the reaction rate is described by r = k[A][B], where k is the reaction rate. k(T) is the reaction rate constant that depends on temperature, where it can be represented in terms of the activation energy:

퐸 − 푎 푘(푇) = 퐴푒 푅푇 where Ea is the activation energy and R is the universal gas constant. This will be revisited in a later portion of the thesis in Chapter 5, Section 5.7, where the reaction temperature was investigated to find the activation energy of the reaction of interest.

The concentrations of the organic compounds (BnOH, BnNH2, Bn2NH and Bn3N) were measured over time by sampling the reaction solution during the run. However, concentrations of NH3 and H2O were not monitored over time and were therefore not included in the data sets. The graphical presentation is shown in Figure 5.2, where the experimental data are shown as data points, and the calculated results are represented by solid curves. The calculated kinetic constants are shown in Table 5.3.

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Concentration (M) BnOH

BnNH2 Bn2NH

Time (hr)

Concentration (M) BnOH

Bn2NH

Bn3N

Time (hr)

Figure 5.2. Reaction profiles of BnOH and NH3 catalysed by 65 wt% Ni/Al2O3/SiO3

for 48 hours (Top) and 1 wt% Au/TiO2 for 4 hours (Bottom). Kinetic profiles were

generated by Berkeley Madonna, where the experimental data are shown as data points

and the calculated results are represented by solid curves.

155

a Table 5.3. Kinetic constants of commercial Ni- and Au- catalysed NH3 reaction A

k3f k3r -1 -1 -1 -1 Entry Catalyst k1f (s ) k1r (s ) k2f (s ) k2r (s ) K1 K2 K3 (s-1) (s-1)

7.1 x 1.8 x 1 Nib 2.3 100 100 5.6 310 3.24 x 10-14 1.0 10-15 10-2

1.5 x 2 Auc 0.92 0 0 1.9 0.18 Irreversible Irreversible 11 103

a Experimental conditions: Batch reactor, 160 °C, 1000 rpm, 0.10 M BnOH in o-xylene

(30 mL), BnOH/NH3/H2O = 1/1/5, t = 0 when T = 160 °C, samples analysed by GC-

FID. b BnOH/Ni = 20, 48 hours, c BnOH/Au = 100, 4 hours

For all reactions, the forward rates kf are larger than the corresponding backward rates,

kr, which suggests that all the forward reactions are favoured. The kf/kr ratio represents

the equilibrium constant K, which gives a better perspective for the overall reaction

equilibrium. For instance, K1 = k1f/k1r.When K is larger than 1, the larger this ratio is

the more favoured it is forwards; when K is smaller than 1, the smaller the value the

more favoured it is backwards. When the reverse kinetic constant kr = 0, the reaction is

considered irreversible. It should be noted that the ratio BnOH/M, where M is any metal,

is different for the two sets of reactions, where BnOH/Au is 5x of BnOH/Ni, which

should be considered when making direct comparisons.

For Ni-catalysed reactions (Table 5.3, entry 1), the first forward reaction was moderate

-1 -15 -1 at k1f = 2.3 s , and the first reverse reaction was much slower at k1r = 7.1 x 10 s . This

14 resulted in an equilibrium constant of K1 = 3.24 x 10 , which suggests the forward

-1 reaction was highly favourable. The second forward reaction was fast at k2f = 100 s ,

-1 but with an equally fast reverse reaction at k2r = 100 s , resulting in an equilibrium

constant of K2 = 1, which suggests the products BnNH2 and Bn2NH were in equilibrium.

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-1 The third forwards reaction was moderate at k3f = 5.6 s , with a much faster reverse

-1 -2 reaction at k3r = 310 s , resulting in an equilibrium constant of K3 = 1.8 x 10 . This suggests that the reverse of the third reaction was much more favoured, and thus no

Bn3N was observed. This is reflected in the product distribution, where BnNH2 was the major product at a selectivity of 85% and Bn2NH at 15%.

For Au-catalysed reactions (Table 5.3, entry 2), the first forward reaction was

-1 irreversible with a small kinetic constant of k1f = 0.92 s , which suggests that the reaction proceeded slowly, but irreversibly. The second reaction was also irreversible,

-1 but much faster with k2f = 1500 s . The third reaction was reversible, with a forward

-1 -1 rate k3f = 1.9 s and reverse rate k3r = 0.18 s . The equilibrium constant of the third reaction, K3 = 11 suggested that the equilibrium position of the third reaction is highly favoured forwards to the products. This was also reflected by the experimental product distribution with a selectivity of 60% towards Bn2NH and 37% towards Bn3N.

When comparing the forward kinetic constants, k1f, k2f and k3f, the forward rates of the

-1 -1 second reaction, where BnNH2 reacted with BnOH (k2f, Ni = 273 s and k2f, Au 1500 s )

-1 -1 were much faster than the rates of the first (k1f, Ni = 2.0 s and k1f Au = 0.92 s ) or third

-1 -1 reaction (k3f, Ni =32 s and k3f, Au = 1.9 s ). This was expected as BnNH2 is a lot more nucleophilic than NH3 and Bn2NH, thus reacting more readily with BnOH.

When comparing the two sets of kinetic constants, the difference in the molar ratios of

BnOH/Ni and BnOH/Au must be considered. The Table below shows the kinetic constants of Ni-catalysed reactions after being multiplied by 0.2 to account for that difference.

157

Table 5.4. Adjusted kinetic constants for commercial Ni- and Au- catalysed reaction A.

-1 -1 -1 -1 -1 -1 Entry Catalyst k1f (s ) k1r (s ) k2f (s ) k2r (s ) k3f (s ) k3r (s )

1 Nib 0.46 1.4e-15 20 20 1.1 62

2 Auc 0.92 0 1500 0 1.9 0.18

After accounting for the molar ratio difference, the first forward kinetic constant for Au, k1f,Au is 2 times that of the first kinetic constant for Ni, k1f,Ni. Furthermore, the second forward kinetic constants k2f,Au is 75 times of k2f,Ni, which suggests that Au was faster than Ni in catalysing the first reaction, and much faster in catalysing the second reaction.

Lastly, k3f,Au is slightly larger than k3f,Ni, which suggests that Au should have catalysed the third reaction faster than Ni. However, the reverse kinetic constant of the third reaction by Ni, k3r,Ni is over 300 times more than k3r,Au, which is reflected in the difference in product distribution, where no Bn3N was observed for the Ni-catalysed reaction and a 37% selectivity to Bn3N was achieved with Au.

As 65 wt%Ni/Al2O3/SiO2 yielded high selectivity towards primary amine, it was selected as the catalyst in the later reactions unless specified otherwise. This catalyst will hereon be referred to as Catalyst 1.

158

5.5. Reactions with Intermediates and Refined Kinetic Model with Ni

In order to understand the reactions steps in greater detail, a series of Ni-catalysed batch reactions were carried out using BnOH with aqueous NH3 as well as the intermediate amines (BnNH2, Bn2NH & Bn3N) with water as starting materials. Each reaction was designed to achieve a constant ratio of Bn/NH3/H2O at 1/2/10. In theory, all reactions should achieve the same thermodynamic equilibrium when the reactions were carried out indefinitely. However, as these reactions were carried out for only 48 hours, full reaction completion were not expected, especially with the inert Bn3N. The reactions were carried out under the conditions of 160 °C for 48 hours with 30 mL of 0.6 M BnOH in o-xylene, 5 mol% Catalyst 1; additional NH3 (aq) and/or H2O were added as required.

The results are shown in Table 5.5.

[푟푒푎푐푡푎푛푡] −[푟푒푎푐푡푎푛푡] The conversion was calculated by 0 × 100%, where the reactant [푟푒푎푐푡푎푛푡]0 was defined as the alcohol or amines used as the starting material.

[푃푟표푑푢푐푡] The selectivity was calculated by 1 × 100% [푃푟표푑푢푐푡]1+[푃푟표푑푢푐푡]2+[푃푟표푑푢푐푡]3

The reaction with BnOH, NH3 and H2O (Table 5.5, entry 1) saw a low BnOH conversion at 27% but a high selectivity to BnNH2 at 97%, with Bn2NH being the only side-product at 3%. This high selectivity matches that of the previous experiment with Ni as the same catalyst. However the conversion was lower than the previous results, possibly due to the addition of water as well as the shorter reaction times.

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Table 5.5. Ni catalysed reactions of product amines with water. Initial concentrations of each substrate are listed in shaded boxes for clarity.

Initial conc. (M)

Final conc. (M) Conversions &

Entry Reaction BnOH NH3 BnNH2 H2O Bn2NH Bn3N Selectivities

Conv.= 27% 0.6 1.2 0 5.4 0 0

1 BnOH+NH3+H2O Sel. (BnNH2 = 97%, 0.12 n/a 0.35 n/a 0.023 0 Bn2NH= 3%)

Conv. = 80% 0 0.6 0.6 6.0 0 0 Sel. (BnOH = 8%, a 2 BnNH2+H2O Bn2NH = 62%, 0.039 n/a 0.12 n/a 0.172 0 PhCHNBn = 30%)

Conv. = 99% 0 0.9 0 6.0 0.3 0

3 Bn2NH+H2O Sel. (BnOH = 90%, 0.60 n/a 0.063 n/a 0.003 0 BnNH2 = 10%)

0 1.0 0 6.0 0 0.2 Conv. = 4%

4 Bn3N+H2O Sel (BnNH2 = 9%, 0 n/a 0.003 n/a 0.015 0.18 Bn2NH = 91%)

Conditions: 160 °C, 1000 rpm, o-xylene (30 mL), [Bn] = 18 mmol,

[Bn]:[NH2]:[OH]:[Ni] = 1:2:10:0.05, 48 hours

The reaction between BnNH2 and H2O saw high conversion at 80% (Table 5.5, entry

2). Selectivity was high towards the secondary amine and imine at 62% and 30%

respectively, compared to 8% for alcohol. This suggests that once the BnOH was

produced via the backwards reaction, it was quickly consumed as the primary amine is

more nucleophilic than NH3, thus producing more of the higher amine. A previously

unobserved product imine, PhCHNBn, was detected for the reaction mixture of

BnNH2+H2O. This imine is an intermediate product of the reaction between BnNH2 and

160

BnOH (Scheme 5.15). The presence of the imine indicates a lack of hydrogen source, as BnOH is not present in the starting reaction mixture.

Scheme 5.15. Hydrogen borrowing cycle of the reaction between BnNH2 and BnOH.

The reaction with Bn2NH and H2O also saw high conversion at 99% (Table 5.4, entry

3). The product distribution suggests that the reaction was favoured backwards. It was highly selective towards BnOH at 90%, with the only side product being BnNH2.

Insignificant reaction happened with the Bn3N and H2O reaction, as evidenced by the low conversions of 4% (Table 5.4, entry 4). This suggests that the reverse reactions at step 2 and 3 were very slow, and the reaction time of 48 hours was not enough.

A previously unobserved product imine, PhCHNBn, was detected for the reaction mixture of BnNH2 and H2O. This imine is an intermediate product of the reaction between BnNH2 and BnOH. As a result, the previous three-reaction model was expanded to four

161 reactions. This expanded process is illustrated in Scheme 5.16 and the differential equations developed for this model are shown in appendix (S5).

Scheme 5.16. A four-reaction kinetic model to describe the reaction between the N- alkylation of NH3 with BnOH (reaction A).

In addition to the six reaction rate constants, as defined previously, the reaction rate constants k4f and k4r are defined as follows:

• k4f = the forward kinetic rate constant for the 4th reaction,

• k4r = the reverse kinetic rate constant for the 4th reaction.

The kinetic constants from the three-reaction model (Set 1A) and four-reaction model

(Set 1B) are shown in Tables 5.6 and 5.7 respectively. A general solver had also been carried out for each model. A general solver (GS) is calculated by considering different data sets and fitting them with a universal set of kinetic values. The data sets may have different initial concentrations but must have the same reaction conditions. The code used to calculate these values can be found in the appendix (S5).

162

Scheme 5.17. A recap of the three-reaction kinetic model used for reaction A.

Table 5.6. Ni-catalysed reactions of product amines with water (set 1A). Kinetic constant values were fitted by using the three-reaction model (Scheme

5.17)

k1f k1r k2f k2r k3f k3r K1 K2 K3

Entry Reactions (s-1) (s-1) (s-1) (s-1) (s-1) (s-1) (s-1) (s-1) (s-1) RMS

-12 10 1 BnOH+NH3+H2O 0.23 2.9 x 10 88 7.9 3 4.4 7.9 x 10 11 0.68 0.12

-2 -2 2 BnNH2+H2O 3.1 2.8 x 10 14.4 8.0 x 10 13 0.26 110 180 50 0.11

-3 -2 -2 -2 -2 -2 3 Bn2NH+H2O 1.6 x 10 0.13 1.0 x 10 2.0 x 10 2.0 x 10 1.8 1.2 x 10 0.50 1.1 x 10 0.24

4 Bn3N+H2O 0.75 0.74 0.75 0.74 0.75 0.74 1.0 1.0 1.0 0.05

5 GS 4.1 x 10-2 2.7 x 10-18 3.7 x 10-2 8.3 x 10-4 1.2 x 10-16 4.5 x 10-2 1.5 x 1016 45 2.7 x 10-15 0.90

Conditions: 160 °C, 1000 rpm, o-xylene (30 mL), [Bn] = 18 mmol, [Bn]:[NH2]:[OH]:[Ni] = 1:2:10:0.05, 48 hours

131

It is expected that the kinetic constant values should not vary too much across the reactions, i.e. k1f should be identical for all four reactions. However, the calculated results showed otherwise. For example, the values of k1f for the four reactions ranged from 1.6 x 10-3 s-1 to 3.1 s-1. However, valuable insight may still be gained when comparing the kinetic values within the same reaction.

Discussion for reaction kinetics using the three-reaction model (Table 5.6)

BnOH+NH3+H2O (Table 5.6, entry 1)

For the reaction involving BnOH, NH3 and H2O (Table 5.6, entry 1), the first forward

-1 reaction was slow at k1f = 0.23 s , and the first reverse reaction was much slower at k1r

-12 -1 = 2.9 x 10 s . This results in an equilibrium constant of K1 = 7.6e10, which suggests that the forward reaction is highly favourable. The second forward reaction was fast at

-1 -1 k2f = 88 s , but with a fast reverse reaction at k2r = 7.9 s , resulting in an equilibrium constant of K2 = 11, which suggests the reaction is favoured forwards. The third forward

-1 -1 reaction is relatively fast at k3f = 3 s , with a moderate reverse reaction at k3r = 4.4 s , resulting in an equilibrium constant of K3 = 0.68. This suggests that the reverse of the third reaction is favoured. This is reflected in the product distribution, where BnNH2 was the major product at a selectivity of 85%, with Bn2NH at 15%.

BnNH2 + H2O (Table 5.6, entry 2)

A similar trend can be observed with the reaction involving BnNH2 and H2O (Table 5.6, entry 2). However, the kfs are much larger than the krs by a factor of 50 to 180, indicating that the forward reactions are highly favoured. The k1f is the smallest when compared to other kf values, again suggesting that the first reaction was the slowest.

132

Bn2NH+H2O (Table 5.6, entry 3)

The reaction with Bn2NH and H2O (Table 5.6, entry 3) tells a different story. The kinetic values for forward reactions are smaller than those of the backward reactions which suggests that the reverse reactions are favoured.

Bn3N+H2O (Table 5.6, entry 4)

The reaction that started with Bn3N and H2O (Table 5.6, entry 4) had very little conversion, thus the calculated forward and backward kinetic constants were very

-1 -1 similar in values, with k1f = k2f = k3f = 0.75 s and k1r = k2r = k3r = 0.74 s . This suggests highly reversible reactions, with the forward reaction being slightly favoured, as reflected in the equilibrium constants where they are all slightly higher than 1.

General solver for set 1A (Table 5.6, entry 5)

The general solver (Table 5,6, entry 5) for this three-reaction model suggested a slow

-2 -1 k1f of 4.2 x 10 s , which agrees with previous findings. The much slower k1r = 2.7 x

-18 -1 10 s suggests that the first reaction highly favoured forwards to BnNH2 formation

16 -2 -1 -4 -1 (K1 = 1.5 x 10 ). The k2f is low at 3.7 x 10 s and the k2r is lower at 8.3 x 10 s , suggesting a slow forward reaction that favoured formation of Bn2NH (K2 = 45). Lastly,

-16 -1 -2 -1 the k3f is very small at 1.2 x 10 s , and the k3r is much larger at 4.5 x 10 s , which suggests the reaction was a slow but highly favoured reverse reaction where Bn3N

-15 slowly converted to Bn2NH (K3 = 2.7 x 10 ).

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Scheme 5.18. A recap of the four-reaction kinetic model used for reaction A.

Table 5.7. Ni-catalysed reactions of product amines with water (Set 1B). Kinetic constant values were fitted by using the four-reaction model (Scheme 5.18).

k1f k1r k2f k2r k3f k3r k4f k4r Entry Reaction K1 K2 K3 K4 RMS (s-1) (s-1) (s-1) (s-1) (s-1) (s-1) (s-1) (s-1)

-4 -2 -2 -3 -5 1 BnOH+NH3+H2O 0.19 3.7 x 10 8.8 x 10 4.6 x 10 0.95 0.44 7.1 x 10 2.2 x 10 510 1.9 2.2 320 0.40

-9 -6 -4 2 BnNH2+H2O 2.8 0.11 12 0.70 0.10 0.30 1.9 x 10 3.0 x 10 25 17 0.33 6.3 x 10 0.09

-2 -2 -2 -2 3 Bn2NH+H2O 0.82 0.67 9.7 x 10 2.0 x 10 0.47 0.90 1.5 x 10 1.1 1.2 4.9 0.52 1.3 x 10 0.55

-4 -5 3 -6 4 Bn3N+H2O 0.49 1.3 0 0.65 2.5 2.9 x 10 0.40 x 10 1.4 0.38 0 8.6 x 10 2.9 x 10 0.046

5 GS 4.2 x 10-2 1.0 x 10-18 3.9 x 10-2 8.0 x 10-4 1.2 x 10-15 0.13 1.3 x 10-18 1.9 x 10-4 4.2 x 1016 49 9.2 x 10-15 6.8 x 10-15 0.87

Conditions: 160 °C, 1000 rpm, 0.60 M BnOH in o-xylene (30 mL), 5 mol% Catalyst 1, [Bn]:[NH2]:[OH] = 1:2:10

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Discussion for reaction kinetics using the four-reaction model (Table 5.7)

BnOH+NH3+H2O (Table 5.7, entry 1)

For the reaction involving BnOH, NH3 and H2O (Table 5.7, entry 1), the first forward

-1 reaction was slow at k1f = 0.19 s , and the first reverse reaction was slower at k1r = 3.7

-4 -1 x 10 s . This resulted in an equilibrium constant of K1 = 510, which suggests that the forward reaction was favourable. The second forward reaction was slow at k2f = 8.87 x

-2 -1 -2 -1 10 s , but with a slightly faster reverse reaction at k2r = 4.6 x 10 s , resulting in an equilibrium constant of K2 = 1.9, which suggests the forward reaction was slightly more favoured. The equilibrium reaction between the imine and amine was moderate at k3f =

-1 -1 0.95 s , with a moderate reverse reaction at k3r = 0.44 s , resulting in an equilibrium constant of K3 = 0.68. This suggests the reverse of the third reaction was favoured. This is reflected in the product distribution, where BnNH2 was the major product at a selectivity of 85%, with Bn2NH at 15%.

BnNH2+H2O (Table 5.7, entry 2)

For the reaction involving BnNH2 and H2O (Table 5.7, entry 2), the intermediate imine,

PhCHNBn was observed and the model was designed to account for it. Therefore, this model should be a better fit than the three-reaction model. This is reflected in the RMS value at 0.09 compared to the previous value of 0.11, which is a slight improvement.

The kinetic constants follow a similar trend from the previous entry, where the first

-1 reaction was moderately fast (k1f = 2.8 s ) with a forward favoured reaction at K1 = 25.

The second reaction was also favoured forwards (K2 = 17) but was four times faster than

-1 the first reaction (k2f = 12 s ). The third reaction, where PhCHNBn was converted into

Bn2NH, was favoured backwards with a K3 of 0.33, which proceeded slowly with k3f

-1 -1 and k3r at 0.1 s and 0.3 s respectively. Lastly, the fourth reaction proceeded very

-9 -1 -6 -1 slowly with k4f = 1.9 x 10 s and k4r = 3.0 x 10 s , which suggests that the reverse

-4 reaction was highly favoured with K4 = 6.3 x 10 .

135

Bn2NH+H2O (Table 5.7, entry 3)

For the reaction involving Bn2NH and H2O (Table 5.7, entry 3), the first reaction was

-1 slow but favoured forwards (k1f = 0.82 s , K1 = 1.2), and the second reaction, which

2 -1 produces the imine, was even slower but more favoured forwards (k2f = 9.7 x 10 s , K2

-1 = 4.9). The third reaction was also slow but favoured backwards (k3f = 0.47 s , K3 =

0.52), Finally, the last reaction was favoured backwards and proceeded at a moderate

-1 -2 speed (k4r = 1.1 s , K4 = 1.3 x 10 ).

Bn3N+H2O (Table 5.7, entry 4)

For the reaction with Bn3N and H2O (Table 5.7, entry 4), the first reaction, again

-1 - proceeded slowly but was favoured backwards in this data set (k1f = 0.49 s , k1r = 1.3 s

1 -1 , K1 = 0.38). The second backwards reaction was considered irreversible, as k2f = 0 s .

-1 This reaction proceeds slowly as k2r = 0.65 s . The third reaction was highly favoured

3 -1 forwards with K3 = 8.6 x 10 , proceeding at a moderate rate of k3f = 2.5 s . The fourth

-6 reaction was highly favoured backwards with K4 = 2.9 x 10 , also proceeding

-1 moderately at k4r = 1.4 s .

General Solver for set 1B (Table 5.7, Entry 5)

The general solver (Table 5.7, Entry 5) for this four-reaction model suggested a slow

-2 -1 -18 -1 k1f of 4.2 x 10 s and a much slower k1r = 1.0 x 10 s suggesting that the first reaction was highly favoured forwards to BnNH2 formation, with the equilibrium constant being

16 -2 -1 very large with K1 = 4.2 x 10 . The k2f was low at 3.9 x 10 s and the k2r was lower at

-4 -1 8.0 x 10 s , suggesting a slow forward reaction favoured to form Bn2NH (K2 = 49)

The third reaction, which was the conversion of the imine to the secondary amine,

-15 showed a highly favoured backwards reaction, with K3 = 9.2 x 10 , which proceeded

-1 -18 -1 slowly at k3r = 0.13 s . Lastly, the k4f was very small at 1.3 x 10 s , and the k3r was much larger at 1.9 x 10-4 s-1, which suggests the reaction was a slow but highly favoured

136 reverse reaction where Bn3N slowly converted to Bn2NH with the equilibrium constant

-15 of K4 = 6.8 x 10 .

Summary

In this section, a series of Ni-catalysed batch reactions were carried out using BnOH with aqueous NH3 as well as the intermediate amines (BnNH2, Bn2NH & Bn3N) with water as starting materials. The secondary imine, PhCHNBn was an unexpected product that was not observed in previous reactions. Thus a 4-reaction model was developed to take the imine into account. However, it was found that the 4-reaction model did not fit the results much better than the 3-reaction model, as evident in the small difference in the RMS of the general solver. Since there was no more PhCHNBn observed in later reactions, the reaction kinetics were analysed using the 3-reaction model.

The production of PhCHNBn brings questions about the hydrogen borrowing cycle and whether it is operating as a full cycle in these reactions. As such, it would be insightful to carry out gas phase analysis on these reactions to detect any production of H2. This will be further discussed in Chapter 7, Section 7.2.

The experimental results did not match the thermodynamic calculations, which may be explained by the very slow formation of the tertiary amine, Bn3N, as calculated by both reaction models.

The following section examines how water affects reaction A. This will be explored by conducting batch experiments with different concentrations of water, and subsequently the removal of water from the initial reaction mixture.

137

5.6. Effect of Water

To understand the extent of how water was affecting the overall reaction equilibrium, further batch reactions of BnOH/NH3/H2O were performed under different ratios, while still using Catalyst 1. These experiments were performed by either:

• adding aqueous NH3 and additional water into the solution (Table 5.7,

entries 1 – 4, set 2A), or

• adding gaseous NH3 into the reactor at different pressures (Table 5.7,

entries 5 -7, set 2B)

It was expected that with a higher water concentration would lower the conversion of

BnOH. This is because the condensation step in the hydrogen borrowing cycle is reversed by water, as shown in Scheme 5.19.

Scheme 5.19. Hydrogen borrowing cycle with emphasis on the condensation step.

In reaction set 2A, where aqueous NH3 was used as a reactant, 28 wt% NH3 solution was used. This solution contained NH3 and H2O at a molar ratio of 1:5 NH3:H2O intrinsically. This was calculated using the density of aqueous NH3 and the molar mass of water. Additional water was added into the reaction mixture as required.

In reaction set 2B, pressurised NH3 gas was used as the NH3 source. The size of the reactor headspace was 10 mL, which allowed a maximum of 5.9 mmol of gaseous NH3 to be present under room temperature and pressure (25 °C, 1 bar). For reactions that

139 required a lower quantity of NH3, the stoichiometry was controlled by lowering the initial pressure (1 – 2 bar NH3). In the reaction with BnOH/NH3 = 1, a smaller volume of BnOH solution (10 mL) was used in order to match the ratio. (Table 5.7, Entry 5), although data over time was not collected due to the small reaction volume used. For reactions that required more NH3 than the reactor headspace allowed, a condensation method was developed.

Introduction of anhydrous NH3: Condensation tube method

A condensation tube was first constructed, which consisted of a closed 0.12 m 1/4” stainless steel tube attached to a needle valve. The picture was shown in Chapter 4,

Section 4.3.1 Batch Reactions, Figure 4.2).

During the condensation of NH3, the tube was submerged in an acetone/ice bath at -10

°C for 20 minutes while attached to NH3 (g) supply. The exterior of the condensation tube was dried to remove excess water or acetone after the cooling process. The tube is then weighed before attaching it to the batch reactor (Figure 5.5). The content was then released into the batch reactor, aided by a heat gun. The final weight after the release of content was then measured again, and the difference was calculated as the injected NH3.

This condensation method allowed accurate measurement of the NH3 being introduced into the batch system. Furthermore, using a smaller sized tube (50 g) is more energy efficient than cooling down the entire batch reactor (> 2000 g).

The condensed NH3 amounted to ~0.70 g (41 mmol), which was the maximum amount of NH3 able to be condensed in the tube. This was confirmed by longer condensation times, which did not increase the total weight of the condenser tube. It was important to carry out all procedures in a fume hood. The entire setup of the batch reactor is shown in Figure 4.1 in Chapter 4, Section 4.3.1 .

The conversions and selectivities from reaction sets 2A and 2B are shown in Table 5.8.

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The thermodynamic predictions indicated 100% conversions under these conditions.

However, the experimental reaction time of 72 hours was not long enough to achieve full conversions, and 72 hours is already a long reaction time. Thus, the reactions were stopped at 72 hours. The high selectivity to BnNH2 was expected from the thermodynamics calculations. In the case where BnOH was used as an excess (Table

5.8, Entry 6), the selectivity to Bn2NH was higher, but no Bn3N was observed, which could be due to the third reaction being a very slow reaction.

Table 5.8. Conversions and selectivities from reaction A catalysed by Catalyst 1, using

a different BnOH:NH3:H2O ratios

Entry BnOH:NH3:H2O (Set 2A) Conversion (%) Selectivities (%)

BnNH2 Bn2NH

1 1:1:5 55 87 13

2 1:1:10 27 93 7

3 1:2:10 52 88 12

4 1:2:20 28 87 13

BnOH:NH3 (Set 2B)

5 b 1:1 86 90 10

6 c 6:1 81 67 33

7 1:3 96 97 3 aConditions: 160 °C, 72 hours, 1000 rpm, 0.60 M BnOH in o-xylene (30 mL), 10 mol%

Catalyst 1; b only 10 mL BnOH solution was used; c 58 hours

No tertiary amines were observed in all the reactions. Although the selectivity to BnNH2 was generally high at 87-93% across all experiments, the conversions were low for

141 reaction set A at 27-52% but much higher for reaction set B at 81-96%. This was expected, and was due to the presence of water in reaction set 2A, which inhibits the forward reactions.

Interestingly, for reaction set 2A, the conversions correlated with the ratio of NH3:H2O with 27 – 28% at 1:10 (Table 5.8, entries 2 and 4) and 52 – 55% at 1:5 (Table 5.8, entries

1 and 3), but were independent of the ratio between NH3:BnOH. (Figure 5.3). The presence of water slowed the forward reaction, as the second step of the hydrogen borrowing cycle involved a condensation reaction.

Figure 5.3. The molar ratio of H2O/NH3 was plotted against conversion.

For reaction set B, selectivity to BnNH2 was comparable at 90% as against 93% in the aqueous systems. Bn2NH selectivity was slightly higher at 10% compared to 7%. It is interesting that the selectivities were similar to systems with aqueous NH3, which proves that the presence of water affects reaction rates.

As predicted, anhydrous NH3 reactions achieved higher conversions when compared with reactions using aqueous NH3 (81 – 96% vs 27 – 52%). Therefore, from here onwards, all reactions were carried out with anhydrous NH3 unless stated otherwise.

142

The online data sets of these experiments were extracted and their kinetic values were calculated by fitting their data sets using the three-reaction model. A general solver (GS) was also calculated for this set of results. The results are shown below in Table 5.9.

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Table 5.9. Kinetic constants from reaction A by Catalyst 1 in batch, using different BnOH:NH3:H2O ratios. The kinetic constant values were fitted by Berkeley

Madonna using the three-reaction model.

BnOH:NH3:H2O Kinetic constant values Equilibrium constants

-1 -1 -1 -1 -1 -1 Entry (Set 2A) k1f (s ) k1r (s ) k2f (s ) k2r (s ) k3f (s ) k3r (s ) K1 K2 K3 RMS

1 1:1:5 3.6 x 10-2 3.0 x 10-16 1.4 1.2 2.1 1.3 1.2 x 1014 1.2 1.6 3.4 x 10-2

2 1:1:10 1.5 x 10-2 2.7 x 10-3 0.23 0.30 0.50 3.1 x 10-2 5.6 0.77 16 6.7 x 10-2

3 1:2:10 1.0 x 10-2 1.6 x 10-3 1.6 1.5 5.5 x 10-16 4.7 x 10-11 6.3 1.1 1.2 x 10-5 4.0 x 10-2

4 1:2:20 2.9 x 10-2 9.9 x 10-4 25 18 2.7 6.4 29 1.4 0.42 4.9 x 10-2

BnOH:NH3

(Set 2B)

5b 1:1 No online data available

6 6:1 0.36 3.4 x 10-5 54.5 7.6 1.6 1.7 x 10-2 1.1 x 104 7.2 94 0.15

7 1:3 0.21 3.5 x 10-2 0.40 0.83 0.91 0.42 6.0 0.48 2.2 7.7 x 10-2

8 GS 1.3 x 10-2 3.5 x 10-17 2.7 x 10-3 0.035 0.052 13.6 3.7 x 1014 7.7 x 10-2 3.8 x 10-3 1.4

aConditions: 160 °C, 72 hours, 1000 rpm, 0.60 M BnOH in o-xylene (30 mL), 10 mol% Catalyst 1, kinetic results are calculated using a 3-reaction model; b10

mL BnOH solution, no online information available

144

Reaction set 2A: BnOH:NH3:H2O reactions (Table 5.9, entries 1 – 4)

BnOH:NH3:H2O = 1:1:5 (Table 5.9, Entry 1)

For the reaction with a BnOH:NH3:H2O ratio of 1:1:5 (Table 5.9, Entry 1), the forward

-2 -1 kinetic value of the first reaction, k1f was slow at 3.6 x 10 s , but the reverse kinetic

-16 -1 value, k1r was much slower at 3.0 x 10 s . This suggests that the equilibrium was highly forward favoured for the first reaction with a very large equilibrium constant

14 value, K1 = 1.2 x 10 . For the second reaction, k2f and k2r were very similar in value at

-1 -1 1.4 s and 1.2 s respectively, and similarly for the third reaction, k3f and k3r were very similar in values at 2.1 s-1 and 1.3 s-1 respectively. The closeness in the kinetic values for the forward and reverse reactions suggests that the equilibrium was slight forward favoured, with the equilibrium constants of K2 = 1.2, K3 = 1.6. Although the equilibrium constants indicated that the reactions should be favoured to Bn3N, no traces of the tertiary amine were observed. This could be because the reaction was very slow, and the reaction time was not long enough.

BnOH:NH3:H2O = 1:1:10 (Table 5.9, Entry 2)

For the reaction with BnOH:NH3:H2O = 1:1:10 (Table 5.9, entry 2), more water was added in the reaction mixture compared to the previous reaction. k1f was slow at 1.5 x

-2 -1 -3 -1 10 s , and k2r was slightly slower at 2.7 x 10 s , which resulted in K1 = 5.6, indicating

-2 -1 -1 a slightly forward favoured reaction. k2f and k2r were slow at 0.23 x 10 s and 0.30 s respectively, with K2 = 0.77, which suggests a slight favoured backwards reaction. k3f

-1 -2 -1 and k3r were also slow at 0.50 s and 3.1 x 10 s respectively, with K3 = 16, which suggests a forward favoured reaction.

BnOH:NH3:H2O = 1:2:10 (Table 5.9, Entry 3)

For the reaction with BnOH:NH3:H2O = 1:2:10 (Table 5.9, Entry 3), twice the amount of aqueous NH3 was added in the reaction mixture compared to Entry 1. The kinetic

145

-2 -1 results show very similar trends to other reactions, with k1f and k1r slow at 1.0 x 10 s

-3 -1 and 1.6 x 10 s respectively, which favoured the forward reaction with K1 = 6.3; k2f

-1 -1 and k2r having very similar values at 1.6 s and 1.5 s respectively, which favoured the forwards reaction slightly with K2 = 1.1; and lastly, k3f and k3r are much slower than in the other experiments, at 5.5 x 10-16 s-1 and 4.7 x 10-11 s-1 respectively, yielding a very

-5 low K3 at 1.2 x 10 , which suggests that the reverse reaction was highly favoured.

BnOH:NH3:H2O = 1:2:20 (Table 5.9, Entry 4)

The last reaction of set 2A, BnOH:NH3:H2O = 1:2:20 (Table 5.9, Entry 4) had the most

NH3 and H2O in the initial setup. The kinetic values follow a similar pattern as before,

-2 -1 -4 -1 with k1f and k1r slow at 2.9 x 10 s and 9.9 x 10 s respectively, which favoured the

-1 forward reaction with K1 = 29; k2f and k2r had values in the same range at 25 s and 18

-1 s respectively, which favoured the forwards reaction slightly with K2 = 1.4; and lastly,

-1 -1 k3f and k3r are higher than the other kinetic rates at 2.7 s and 6.4 s respectively suggesting faster reactions. The third reaction yielded a low K3 at 0.42, which suggests the reverse reaction was moderately favoured.

The kinetic constants of reaction set 2A followed a similar pattern across the four reactions, with a slow but forward first reaction and a fast second reaction with an equilibrium constant very close to 1. However, the third reaction did not follow any pattern, with a forward rate ranging from very slow to moderate, and a range of equilibrium constants at 1.2 x 10-16 s-1 to 16 s-1.

Reaction set 2B: BnOH:NH3 reactions (Table 5.9, entries 5-7)

For reaction set 2B, anhydrous NH3 was used instead of aqueous NH3 in the experiments. Even though there was no water in the initial experimental setup, water was generated over time in the reactions. It was therefore important to keep the reaction model as reversible reactions.

146

BnOH:NH3 = 1:1 (Table 5.9, Entry 5)

As the reaction of BnOH:NH3 = 1:1 (Table 5.9, Entry 5) had a lower reaction solution volume of 10 mL, no data was collected over time, and the kinetics could not be modelled.

BnOH:NH3 = 6:1 (Table 5.9, Entry 6)

For the reaction of BnOH:NH3 = 6:1 (Table 5.9, Entry 6), the k1f is larger than that of

-1 the reactions from set 2A, giving 0.36 s which was 10 times that of Entry 1. The k1r

-5 -1 4 was small at 3.0 x 10 s , yielding a large equilibrium constant, K1 of 1.1 x 10 , which

-1 -1 suggests a highly favoured forward equilibrium. The k2f and k2r are 7.6 s and 1.6 s respectively, suggesting a faster second reaction than the first, yielding K2 at 7.2, which

-1 suggests a moderately favoured forward reaction. The k3f and k3r are at 1.6 s and 1.7 x

-2 -1 10 s respectively, yielding K3 at 94, suggesting a forward favoured equilibrium.

BnOH:NH3 = 1:3 (Table 5.9, Entry 7)

For the reaction of BnOH:NH3 = 1:3 (Table 5.9, Entry 7), k1f = 0.21 and k1r is smaller

-2 -1 at 3.5 x 10 s , yielding an equilibrium constant, K1 of 6.0 for a slightly forward

-1 -1 favoured equilibrium. The k2f and k2r are 0.40 s and 0.83 s respectively, which yields

K2 at 0.48, which suggests a slightly backward favoured reaction. The k3f and k3r are at

-1 -1 -2 0.42 s and 6.0 s respectively, yielding K3 at 7.7 x 10 , suggesting a reverse reaction favoured equilibrium.

The RMSs of the individual fits are very low in the range of 3.4 x 10-2 to 0.15, which indicate the curves were fitted very well.

General Solver for set 2 (Table 5.9, Entry 8)

The GS (Table 5.9, Entry 8) showed similar trends to previous findings. The first

-2 -1 forward reaction was slow at k1f = 1.3 x 10 s with a very slow reverse reaction at k1r

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-17 -1 14 = 3.5 x 10 s . The equilibrium constant K1 = 3.7 x 10 suggests a highly favoured

-3 -1 forward reaction. The k2f is 2.7 x 10 s , which suggests that the second forward

-2 -1 reaction proceeded about 5 times slower than the first reaction. The k2r is 3.5 x 10 s ,

-2 and the equilibrium constant K2 is 7.7 x 10 , which suggests a reverse favoured reaction.

-2 -1 -1 Lastly, the k3f is also small at 5.2 x 10 s , and the k3r is relatively large at 13.6 s . The

-3 equilibrium constant K3 is 3.8 x 10 and the backwards reaction was highly favoured.

General solver comparisons

In order to compare the GS of this set of reactions with that of the reaction set 1A, the

conditions had to be uniform. This set of experiments used 10 mol% of Catalyst 1, which

was double that of the previous set (Table 5.10), which used 5 mol% of Catalyst 1. To

account for this, the kinetic constants of this set of reactions were halved to account for

the difference. The modified set of kinetic constants are shown in Table 5.9.

Table 5.10. General solvers calculated from reaction sets 1A and 2.

Kinetic constants

Entry General Solver k1f k1r k2f k2r k3f k3r

1 Set 1A 4.1 x 10-2 2.7 x 10-18 3.7 x 10-2 8.3 x 10-4 1.2 x 10-16 4.5 x 10-2

2* Set 2 6.5 x 10-3 1.8 x 10-17 1.4 x 10-3 1.8 x 10-2 2.6 x 10-2 6.8

*The kinetic constants have been halved to account for the difference in mol% catalyst

Even though the results were acquired from two different sets of results, the kinetic

values calculated by the general solver match well, which suggests that the model

describes the reaction well. The k1r and k2r are very close in values, with very small

factors of difference of 7 to 20. However, the kinetic rate constants for the third reaction,

k3s do not match well. This could be because there was no Bn3N detected in the system

and thus the model was unable to estimate accurate values of k3r.

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Summary

To summarise, the key rates in the kinetic model for reaction A using Catalyst 1 were the first and second reactions. The first reaction was usually quite slow but was highly forward favoured, and the second reaction usually faster, but had an equilibrium constant close to 1. In addition, Bn3N was not detected in these reactions. Therefore, it can be assumed that the third reaction did not proceed, or the reaction time was not long enough to produce significant amounts of the tertiary amine. Thus, the reaction kinetics model can be slimmed down to a two-reaction model.

The following section, Section 5.7, describes a temperature study for the batch reactions that was carried out at 80, 120 and 160 °C. The reaction kinetics of these reactions are also studied using a two-reaction model.

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5.7. Temperature Studies

In this section, the temperature effects on reaction A will be investigated, as well as the estimation of the activation energy (Ea) and the pre-exponential or the frequency factor, based on the first step of reaction A. Higher temperatures should yield faster reaction rates. This is because the higher the temperature, the more energy the molecules possess, and they would collide more often with a higher rate of successful reactions.

Furthermore, performing this reaction at different temperatures can provide valuable insight into the activation energy.

Three batches of reaction A were conducted at 80, 120 and 160 °C. These were batch reactions carried out under the following conditions: condensed NH3 was used at

NH3/BnOH = 7, 30 mL of 0.6 M BnOH solution in o-xylene. The mixture was reacted for 24 hours using 20 mol% of Catalyst 1 at 80, 120 or 160 °C. The conversions and selectivities are shown in Table 5.11. The kinetic values were calculated by Berkeley

Madonna using a two-reaction model, as shown in Scheme 5.20, where only the first two reactions to BnNH2 and Bn2NH were described. This model was used instead of the three-reaction or four-reaction models because no PhCHNBn or Bn3N was detected.

Scheme 5.20. The two-reaction model used in calculating the kinetic constants for the temperature studies of reaction A.

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Table 5.11. Temperature study of reaction A using Catalyst 1. Entry Temperature Conversion Selectivity to k1f k1r k2f k2r RMS

(°C) (%) BnNH2 (%)

1 80 4 100 5.0 x 10-4 0 0 0 0.050

2 120 55 93 0.013 0 0.17 9.9 0.86

3 160 96 97 0.21 0.031 0.26 0.58 0.072

Conditions: 24 hours, 160 °C, 1000 rpm, 0.60 M BnOH in o-xylene (30 mL), 20 mol%

Catalyst 1, NH3/BnOH = 7, 4.2 bar at 25 °C

At 80 °C, there was only 4% conversion to BnNH2 after 24 hours. (Table 5.11, Entry 1)

At 120 °C, conversions increased to 38%. (Table 5.11, Entry 2) At 160 °C, there was

high conversion of 96% (Table 5.11, Entry 3), matching the conversion of 93% reported

by Shimizu et al. [122]. This increase in conversion as temperature increases was

expected, as the rate of reaction is directly proportional to the reaction temperature.

The selectivity towards BnNH2 at 80 to 160 °C remained high at 93 – 100%, with the

only side-product being Bn2NH at 0 – 7%. The selectivities towards BnNH2 were higher

than the 87% reported by Shimizu et al. [122] This could be due to the higher excess of

NH3 used in these reactions at NH3/BnOH = 7, as facilitated by the condensation

method, compared to only NH3/BnOH = 2.2 used by Shimizu et al..

The kinetic values k1f and k2f increased as temperature increased. k1f, 80°C was very low

at 5.0 x 10-4 s-1 (Table 5.11, Entry 1), which increased to 0.013 s-1 at 120 °C (Table 5.11,

-1 -1 Entry 2) and 0.21 s at 160 °C (Table 5.11, Entry 3). k2f, 120°C was 0.17 s , increasing

-1 only slightly to 0.26 at 160 °C. As k1r, 80 °C and k1r, 120 °C were 0 s , the first reactions at

80 °C and 120 °C were considered irreversible. Furthermore, the second reaction at 80

-1 -1 °C did not occur as k2f, 80 °C = 0 s and k2r, 80 °C = 0 s . The second reaction for 120 °C

-1 -1 and 160 °C were faster than the first reaction, with k2f, 120 °C = 0.17 s , k2r, 120 °C = 9.9 s ,

-1 -1 k2f, 160 °C = 0.26 s and k2r, 160 °C = 0.58 s . However, the reverse reactions were more

favoured for both temperatures.

151

Arrhenius Plot

The k1f of BnOH at 80 – 160 °C were selected to determine the activation energy using the Arrhenius equation:

퐸 − 푎 푘 = 퐴푒 푅푇 (Equation 5.8)

Where k is the kinetic constant in s-1, A is the pre-exponential factor or the frequency

-1 -1 factor in s , Ea is the activation energy in kJmol , R is the universal gas constant in

-1 -1 kJmol K , and T is the reaction temperature in K. k1f was selected as the kinetic constant to calculated A and Ea for the first reaction,. Equation 5.8 can rearranged to

Equation 5.9, whereby a linear graph can be plotted.

퐸 1 ln 푘 = ln 퐴 − 푎 ( ) (Equation 5.9) 푅 푇

The gradient of the graph can be used to find Ea of the reaction, whereas the y-intercept can be used to find A. The plot is shown below as Figure 5.4 and the values of Ea and

-1 -1 A are calculated as Ea = 98 kJmol , A = 9.7e11 s .

Figure 5.4. Arrhenius plot using data from anhydrous NH3 alkylation with benzyl alcohol at 80 – 160 °C.

152

There is currently no work in the literature that describes this activation energy of NH3 monoalkylation with benzyl alcohol. However, Spahlinger and Jackson [157] has computed the activation energies for Menshutkin alkylation of ammonia (Scheme 5.21), in which ammonia displaces several leaving groups via an SN2 mechanism. For example, methyl sulfonate has an Ea of 108 kJ/mol, methyl ester 163 kJ/mol and methyl amide 244 kJ/mol without the aid of catalysts. This suggests that the activation energy for the N-alkylation of NH3 with BnOH to produce BnNH2 is lowered by the catalyst.

Scheme 5.21. Menshutkin alkylation of ammonia, where R is the leaving group and X is the halide.

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5.8. Batch Reactions with Different Alcohols

Different alcohols were tested with the batch system using 36 wt% Ni catalyst extrudites

(Catalyst 2), and the results are shown in Table 5.12. This new catalyst was used for

better comparison with results in the next chapter, because it had a larger size which can

be used in flow reactors. The scope of alcohols included aromatic alcohols, namely

benzyl (entries 1 – 3), 2- & 3- methoxybenzyl (entries 4 & 5) and 2- & 3- picolyl

alcohols (entries 6 & 7), and alkyl alcohols, which are 2-octanol (Entry 8) and 2-

phenylethanol (Entry 9). Additionally, two alkyl diols, 1,5-pentadiol and 1,6-hexadiol

were tested (entries 10 & 11). Conditions were anhydrous NH3/BnOH = 7 with 0.1 M

alcohol in o-xylene, 72 hours, 10 mol% Catalyst 2 in the batch reactor.

a Table 5.12. Results of N-alkylation of NH3 with different alcohols using Catalyst 2 .

Entry Starting material Product Conversion (%) Sel. to primary

amine (%)

1 *(92) 96 *(96) 97 2 b 95 96 3 c 86 99 4 26 99

5 24 99

6 18 99

7 41 96

8 10 0

9 8 55

10 12 99

11 12 99 a (12 hours) 72 hours, 1000 rpm, NH3/ROH = 7, 160 °C, 30 mL of 0.1 M ROH in o-

xylene, ROH/Ni = 10, 10 mol% Catalyst 2; b repeated use of catalyst after one run; c

repeated use of catalyst after two runs.

154

Conversions to primary amines were high for aromatic alcohols, ranging from 96 – 99%.

Among them, BnOH had the highest conversion at 96% after 72 hours (Table 5.12,

Entry 1). After repeated uses of the same catalyst with BnOH, there was a decrease in the conversion from 96% to 86%, but selectivity to primary amine remained high at 96

– 99% (Table 5.12, Entry 2 & 3). The decrease in conversion might be due to carbonaceous deposition, which will be discussed in Chapter 6, Section 6.8, Catalyst

Deactivation.

2-Methoxybenzyl alcohol has an electron donating effect by the electron-rich methoxy group at the ortho- position, and thus is expected to give a higher conversion than 3- methoxybenzyl alcohol, which has the methoxy group at the meta- position (Table 5.12, entries 4 & 5). Interestingly the conversions were similar at 26% and 24% respectively.

This could be due to the steric hindrance of the 2-methoxy group, which decreases the accessibility of the aldehyde carbon, compared to the unhindered 3-methoxy group

(Scheme 5.22). Selectivities for both methoxybenzyl amines were high at 99% and no side products were observed.

Scheme 5.22. Steric effects of a 2-methoxyl benzaldehyde (left) compared to a 3- methoxyl benzaldehyde during the condensation reaction with NH3.

Picolyl alcohols were not previously demonstrated with nickel catalyst (Table 5.12, entries 6 & 7). 2-picolyl alcohol had a conversion at 18%, which was much lower than that of 3-picolyl alcohol at 41%. This can be attributed to the mesomeric effects of the pyridine, which had a positive impact on 3-picolyl at the meso- position, but not present on 2-picolyl aldehydes at the ortho- position. Cartier et al. discussed the mesomeric

155 effects of different positions on pyridines and found that they are similar to benzene

[158] (Scheme. 5.23). Green eluent was observed from samples of 2-picolyl amine, which suggests leaching of Ni from the catalyst. However, ICP was carried out with the eluent, although it was below the detection limit (0.2 ppm).

Scheme 5.23. Resonance effect of a 3-picolyl benzaldehyde (left) compared to the absence of the effect in a 2-picolyl benzaldehyde.

Alkyl alcohols conversions were much lower than those of aromatic alcohols (Table

5.9, Entries 8 & 9). This could be due to the lack of activating benzyl group for the carbonyl to react with NH3. 2-phenylethanol had a selectivity to the amine at 55%, which was much lower than that of aromatic alcohols. There was also no 2-octylamine detected, and 2-octanone was the only product. This suggests that the reaction had stopped at the initial oxidation step and had not proceeded further (Scheme 5.24). This may be attributed to ketones being much less reactive than aldehydes, thus the imine takes longer to form.

Scheme 5.24. Incomplete hydrogen borrowing cycle of 2-octanol, resulting in ketones but no imine or target amines.

156

Both the diols (entries 10 & 11) had slightly higher conversions than 2-phenylethanol at 12%. Interestingly, they also gave 99% selectivity to the diamines, with no side- products observed (as confirmed by GC-MS). However, it should be noted that the diols are highly polar and the mixture with o-xylene was not fully homogeneous at room temperature. This could result in better mixing with NH3 and thus the high selectivities.

Selectivities for aromatic alcohols were generally high at 96 – 99%. However, alkyl alcohols had a much more diverse range of selectivities, with the primary alkyl alcohols at 55 – 99%, and the secondary alcohol at 0% to the amine with only the ketone as the product. This is particularly interesting as Shimizu et al. reported high yield with this secondary alcohol, using a synthesised Ni/CaSiO3 which achieved 86% yield at 20 hours with 1 mol% Ni. However, their system took longer for benzyl alcohol, at only 70% yield with 5 mol% Ni [123]. They suggested that Ni metal particles are the catalytically important species, with the average particle size at 3.0 nm compared to 4.2 nm, as estimated by XRD and TEM (Appendix, P. 228 - 230).

5.9. Comparisons with Different Primary Amine Production Methods

As discussed in Chapter 2 (Literature Review) there are currently six methods in primary amine production. These methods have different reaction mechanisms, so the conditions varied greatly. Yang et al. [84] reported the first example of a one-pot, two- step hydroamination of olefins using a homogeneous Pd/Ir dual metal tandem catalyst system. A turnover number (TON) for the most efficient example was 94

157

[mmolamine]/[mmol catalyst] over 11 hours of total reaction time. This value is higher than the TON of this work with 9.6 [mmol BnNH2]/[mmol Ni] over 72 hours.

Scheme 5.25. Formal one-pot, two-step hydroamination of olefins using a homogeneous Pd/Ir dual metal tandem catalyst system [84].

In the hydroaminomethylation of limonene by Behr et al. [92], the TON of their system was 1 [mmol amine]/[mmol Rh] in 6 hours, which is lower than the TON of this work with 9.6 [mmol BnNH2]/[mmol Ni] over 72 hours.

Scheme 5.26. Hydroaminomethylation of limonene by Behr et al. [92].

The reaction time of the N-alkylation of NH3 with BnOH using Catalyst 1 is much longer than the examples in literature, which suggests a very slow reaction compared to other reactions.

158

5.10. Conclusion

In this chapter, N-alkylation of NH3 with alcohols (reaction A) in batch reactors has been extensively explored.

Firstly, the thermodynamic outcomes of reaction A with different water concentrations are explored under different pressures and temperatures, which had not been previously done. The stoichiometry of the reactants played a major role in product selectivity, which can be controlled by increasing either BnOH or NH3. When water was used in excess, it had a minute effect on the overall outcome.

Preliminary batch reactions were carried out using commercial catalysts. It was found that the Ni catalysts had high selectivities to the primary amine, but were generally slower than the noble metal catalysts, which were highly selective towards the secondary amine. Among them, Au or Ru had not been previously demonstrated to catalyse reaction A.

A three-reaction kinetic model was then built for reaction A catalysed by Ni and Au – and the kinetic values of the forward and backward reactions were determined using

Berkeley Madonna. For both catalysts, the forward rate of the first reaction (k1f) was slower than the other two reactions (k2f & k3f). It was also found that the forward reaction rates of the second and third reactions (k2f and k3f) of Ni were much smaller than that of

Au, which confirmed the difference in product selectivities. The thermodynamic calculations showed that the favourable product under the reaction conditions is BnNH2.

However, the experimental values demonstrated otherwise, with Ni being selective to

BnNH2 and Au selective to Bn2NH. This confirmed that the reactions are controlled by reaction kinetics.

Batch reactions with intermediate amines and H2O were then carried out and their reaction kinetics were analysed. A four-reaction model was developed to account for the intermediate imine. However, it was found that the four-reaction process was not

159 superior over the three-reaction process as the curves were fitted only marginally better, which does not compensate for the cost of the model upgrade. Futhermore, no intermediate imines were observed in later reactions.

Reactions A with different H2O and NH3 were then performed and the kinetics analysed.

It was proven that water has an adverse effect to the forward reaction, and aqueous NH3 was not used in the later reactions. In light of this, an anhydrous NH3 condensation method was developed, which allowed for an actual measurement of NH3 introduced into the batch system, and a higher amount of NH3 than otherwise allowed.

Temperature studies were performed over 80 – 160 °C and reaction kinetics were also studied using a two-reaction model. Higher temperature yielded higher conversions and

-1 -1 faster rates as expected. The Ea and A factor were 98 kJmol and 9.7e11 s , respectively. The Ea of this reaction is high, while the A factor is slow, when compared to a common SN2 alkylation reaction.

To broaden the scope of reactants, a number of alkyl alcohols underwent NH3 alkylation reactions with varying degrees of success, with some conversions ranging from 12 –

96% and selectivities to primary amines at 96 – 99%. However, two of the alcohols had low conversions and selectivies to the corresponding aldehyde or ketone. These batch reaction results were compared with other batch processes, as mentioned previously in the literature review. The results in this work achieved higher selectivities, but with much longer reaction times.

Using information gained from the batch reactions, the N-alkylation of ammonia with alcohols was performed in a flow reactor. The following chapter will be discussing these reactions in detail.

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Chapter 6: Flow Reactions

6.1. Introduction

As previously discussed, flow chemistry has gained increasing attention in recent years for the economic and environmental reasons [18]. Nowadays, it is more commonly used in laboratory settings for reasons such as

• safety,

• more efficient mass and heat transfer, and

• steady state operation for higher consistency [33].

Reactants Products

Scheme 6.1. A simplified scheme of a generic plug flow reactor.

The commercial 65 wt% Ni/Al2O3/SiO2 catalyst demonstrated great conversions and selectivity to the primary amine in batch systems. However, there was an issue with the small catalyst size. The unmodified form of this catalyst is a fine powder (38 - 90 μm) and would not be appropriate for employment in a packed bed due to the high back pressure as well as potential catalyst run-off, which could lead to reactor shut-down and devastating damage to equipment [135].

In an attempt to replicate the high Ni loading catalyst with larger particle size, an extensive search on the production method of the commercial 65 wt% Ni/Al2O3/SiO2 was carried out, which led to a catalyst synthesis patent by Oudejans et al. [136]. Using this co-precipitation method, catalysts of 43 wt% Ni/Al2O3/SiO2 and 85 wt%

Ni/Al2O3/SiO2 were synthesised. These catalysts were tested with reaction A in batch and discussed in detail in Chapter 4.3. In summary, these synthesised catalysts had lower conversions and selectivities to BnNH2 compared to Catalyst 1. Despite their higher surface area, both catalysts had only NiO on the surface and thus was not active

161 in the reaction. An in-situ reduction method was not available because the batch reactor cannot be heated up to 500 °C. An ex-situ catalyst reduction was used, but the activity remained low. As this synthetic method was not successful, an alternative method had to be used to increase the catalyst size.

This was achieved by an extrusions method, which was modified from the works by

Melero et al. [138]. They agglomerated a zeolite with bentonite clay to form macroscopic structured catalyst particles to be used in continuous epoxidation processes on a fixed bed reactor. The extrusions were made with an active catalyst, an inorganic agglomerant clay and an organic additive as a binder. This catalyst was confirmed to contain 36 wt% Ni by induced coupled plasma and was tested with batch reactions with very good performance.

There are many advantages of flow reactions over batch reactions in this project. Flow reactor enables higher pressures in a safe environment, as only a small amount of reaction mixture is heated at one time. This is particularly important with NH3 as a reactant due to its toxicity and volatility, and potentially allows for a higher stoichiometric ratio of NH3/BnOH under liquid phase compared to a batch setup, which may lead to higher selectivity to primary amines. The small reaction space allows for a high catalyst to reactant ratio under super-heated conditions, which leads to shorter reaction times compared to batch reactions. Flow reactions are also more consistent than batch reactions, as the product distribution at steady state operation is very uniform, compared to batch reactions where the products can differ between batches. Moreover, by adjusting the flow rates during the reaction, the residence time can be changed easily, compared to a batch reaction where the entire reaction must be prepared from scratch.

This leads to much faster reaction condition improvements as other factors such as temperature and reactant molar ratio can be adjusted quickly.

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In this chapter the following points will be discussed:

• Preliminary flow reactions with mystery “intermediate”

• How different conditions affected reaction outcome, this includes

o Temperature

o Catalyst size

o Reactant molar ratios

o Flow rates, and

o Initial alcohol concentration

• The scope of alcohols, and

• TON over extended periods of time

6.2. System Characterisation

The physical properties of a reaction mixture must be well understood, particularly regarding whether the reaction happens in a gas or liquid phase. In the present study, the reaction conditions were defined in batch and it is important to confirm a liquid phase reaction during flow conditions. This is because if gas pockets were present in the catalyst bed, issues such as partial stoichiometric imbalance, mass transport and gas evolution might occur. These can lead to unexpected reaction outcome and/or health and safety concerns.

Several simulations were carried out using ASPEN and COFE, and the results discussed in the experimental chapter (Chapter 4, Section 4.6). It was confirmed that under the operating conditions of 160 °C and 60 bar, the system operates in the liquid phase. As shown in Figure 4.11.

Furthermore, the calculation of the Reynold number (Chapter 4, Section 4.3.2) determined that the reaction flow is considered laminar, as Re = 4.3 is < 200.

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6.2.1. Flow Reactor Considerations and Specifications

Commercial flow reactors (such as the X-CubeTM by ThalesNano) were considered for the employment of Reaction A under flow conditions. However, the use of ammonia limits the scope of such devices, as some common plastics such as Nylon (type 6/6) and

Viton are susceptible to ammonia corrosion. Furthermore, the use of a customised flow reactor can allow for a convenient setup for an in-situ catalyst pre-reduction procedure are proved to be crucial in Chapter 6.3.

Using reaction conditions determined as described in Chapter 5, the flow reactor was built with the following aspects in mind:

• high pressure (up to 70 bar) – the back-pressure regulator must regulate from 0-

80 bar, with the pressure relief valve set at 80 bar for pressure spikes;

• temperature range (up to 500 °C) – the heater and temperature controller must

regulate up to 500 °C, where 120 – 200 °C is the operating temperature for

Reaction A, and 500 °C is the reduction temperature;

• packed-bed reactor – the reactor cartridge should be changed regularly with

ease, and a filter with a <250 µm mesh should be installed to prevent any runoff

catalyst particles downstream, as this can potentially damage the back pressure

regulator and cause blockages in the reactor;

• liquefied ammonia delivery – the reactor must deliver condensed NH3 and

withstand NH3 corrosion. All parts used in the reactor should be resistant to NH3

corrosion. Common sealing plastics such as Viton O-rings were avoided;

• safety in the delivery of chemicals - liquefied NH3 is used in the reactor, which

can vapourise immediately at atmospheric room pressure. Therefore, the flow

reactor should be contained in a fume hood. Extra care must be taken to ensure

there are no leakages before running any experiments;

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• removal of NH3 from sample mixture downstream – NH3 must be released

before reaching the sample collection port. This can be achieved by using a gas-

liquid separator. The gas outlet is connected to an acid trap (1 M HCl (aq)

solution), and the liquid sample is collected and sampled using Gas

Chromatography with Flame Ionisation Detector (GC-FID).

A flow reactor was constructed to fulfil these requirements. The reactor scheme is shown previously (Chapter 4, Scheme 4.1). The component details such as their brand, function and capacity are shown in Chapter 4, Section 4.3.2.

In addition to the flow reactor, the NH3 condensation cylinder was constructed and shown below in Figure 6.1. This condensation unit contains two on/off valves that control the inlet and outlet. 1/16” tubing was used inside the cylinder, with the outlet tubing longer than the inlet tubing by 0.10 m. This is because the outlet tubing was submerged in liquid NH3 once condensed which makes liquid extraction easier.

Figure 6.1. Schematic of the NH3 condensation cylinder.

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6.3. Preliminary Flow Reactions

In the preliminary reactions, the NH3 condenser cylinder was first cooled in N2 (l), then charged with NH3 (g) for 2 hours. The cylinder is closed after the charging time until required for liquid NH3 flow. Meanwhile, reactor cartridge (1/4” SS316 tubing, 0.28 m) was packed with 2.0 g of catalyst extrudites (Catalyst 2), suspended with glass wool.

Then the reactor cartridge was installed into the heating cylinders with a heating jacket, and the reactor was filled with o-xylene at a flow rate of 2.0 mL/min using a High

Performace Liquid Chromatography (HPLC) pump. It is important to ensure there are no leakages at this point. The heating unit was then turned on to 160 °C and the pressure is regulated at 60 bar. Once the reaction conditions were met, the pump feed was then switched to 0.6 M BnOH solution in o-xylene and the NH3 condenser cylinder is opened and NH3 (l) was introduced into the reactor at the appropriate flow rates. For example, a combined flow rate of 0.2 mL/min with an NH3/BnOH = 7, where the BnOH solution is 0.1 M would require flow rates of NH3 (l) = 0.018 mL/min and BnOH in o-xylene =

0.182 mL/min. This was calculated according to the molecular weight (MW = 17 g/mol)

-3 and density of NH3 (l) (ϱ = 609 kg m ) [139]. The full calculation of these combined flow rates with different reaction stoichiometry can be found in Chapter 4, Section 4.4,

Table 4.4.

Once the pressure had been re-adjusted for the lower flow rate, the sample collection begins after 30 mL of solution/solvent has been collected. Samples are analysed by GC-

FID and reaction steady-state is achieved after 15 x τ and maintained for 3 x τ, where τ is the residence time. All results are under steady-state conditions unless stated otherwise.

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6.3.1. Unexpected Intermediate: Benzonitrile

During a preliminary experiment, an unexpected, unidentified compound was observed in the product solution. It had a retention time of 3.3 min on GC-FID (column = HP5), which is identical to that of benzaldehyde (PhCHO). This compound appeared as an intermediate, and had a production of 0.08 M at 1 hour, which decreased steadily down to 0 M after 4 hours. This behaviour is not previously seen in batch reactions, and

PhCHO should react with NH3 to produce benzyl amine (BnNH2). Therefore, further identification of the mystery compound was required. Gas Chromatography-Mass

Spectrometry (GC-MS) then identified the compound as benzonitrile, PhCN m/z = 103, where benzaldehyde, PhCHO m/z = 106. A screenshot of the gas chromatograph of one such sample is shown in Figure 6.2, where each compound is labelled.

Solvent

BnOH Internal PhCHNBn PhCN BnNH2 Standard

Figure 6.2. Gas chromatograph of a sample from a flow reaction containing PhCN

The result of this flow experiment over time is shown in Figure 6.3. The reaction conditions were high at 0.6 M BnOH in o-xylene, at 160 °C, 0.2 mL/min, with 2.0 g of

Catalyst 2 and NH3/BnOH at a ratio of 7.

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Figure 6.3. Data collected from flow Reaction A. Benzonitrile (PhCN, black squares) is observed as the major product initially. Conditions: initial BnOH = 0.6 M, 160 °C,

0.2 mL/min, NH3/BnOH = 7, 2.0 g Catalyst 2.

BnNH2 concentration increased from 0 M to 0.15 M after 2 hours and remained steady.

Production of PhCN was observed to peak from 0 M to 0.08 M at 1 hour, which slowly decreased to 0 M after 4 hours, which acts like an intermediate. N-benzyliene- benzylamine (PhCHNBn), the intermediate imine to the secondary amine, Bn2NH, was also observed, increasing from 0 M at 0 hours to 0.05 M after 4 hours. PhCHNBn is usually produced when there is insufficient hydrogen in the system. The production of the nitrile and imine can occur when the catalyst was not activated by hydrogen, as suggested in Eller et al. [159]. They reported that in the absence of hydrogen, imines, enamines, and even nitriles are formed. The continuous hydrogen flow also maintains the catalyst activity by removing carbonaceous deposits and metal carbides or nitrides and prevents disproportionation of amine products. The mechanism of benzonitrile formation is shown in Scheme 6.2.

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Scheme 6.2. Possible pathway of benzonitrile formation from the hydrogen borrowing cycle, when the catalyst was not fully reduced.

With insufficient hydrogen in the system, the reaction did not carry throughout the full hydrogen-borrowing cycle. Instead of reducing benzimine (PhCHNH) to BnNH2,

PhCHNH was oxidised to PhCN. This was caused by high concentration of NiO on the catalyst surface at the beginning of the reaction. As [NiO] decreases and the amount of

[Ni] increases over time, the reaction pathway shifts from imine oxidation to reduction, thus producing more amine. As a result, PhCHNH has been consumed as a reductant over the first four hours of the reaction.

The area under the curves of PhCN and PhCHNBn from Figure 6.2 can be integrated to deduce that 0.96 mmol and 0.48 mmol respectively were produced over the 4 hours.

This implies that at least 1.44 mmol of NiO was initially present in the system.

Subsequent reactions were carried out after catalyst reduction. Prior to the flow reaction setup illustrated above, the catalyst bed was treated under 40 mL/min of 10% H2/N2 at

500 °C for 2 hours, then cooled under 30 mL/min N2 to room temperature (25 °C) for 2 hours prior to any flow experiments. The flow diagram of this set up is shown in Scheme

4.5. The detailed catalyst reduction procedure was described in the Chapter 4, Section

4.3.2. This procedure was modified from the method described by Bartholomew and

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Farrauto [160], where they reduced their Ni/Al2O3 catalyst at 500 °C to achieve a 100% reduction of their Ni catalyst where the nickel surface area marginally changed.

The reaction outcome changed after this reduction method. The results from the first 4 hours of the flow reaction are shown in Figure 6.4. For this reaction, the BnOH concentration was lowered from 0.6 M to 0.1 M, as well as a lower flow rate from 0.2 mL/min to 0.06 mL/min. The other reaction conditions were kept at 160 °C, 60 bar,

NH3/BnOH = 7.

Figure 6.4. Reaction profile of flow reactions using reduced Catalyst 2. Conditions were

0.06 mL/min, 0.1 M BnOH in o-xylene, 160 °C, 60 bar, NH3/BnOH = 7, 2.0 Catalyst 2.

The graph is showing results from the first 4 hours of the reaction.

When compared with the previous graph (Figure 6.3), the production of PhCN has been suppressed in the first four hours, with only BnNH2 production which reached up to

0.05 M at four hours. This reaction was extended to 14 hours long to achieve a steady state, as shown in Figure 6.5.

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The initial rise of concentrations was due to the breakthrough time, where 6 hours was required to fill the entire flow reactor with the reaction solution, and a further 2 hours to reach steady state at 8 hours. During steady state, 96% of BnOH was converted, and selectivity to BnNH2 was at 97% and PhCN was only at 3%. A significant improvement from the unreduced catalyst.

There is a clear difference between the reaction profiles of unreduced (Figure 6.3) and reduced catalysts (Figure 6.5). With the unreduced catalyst, PhCN was observed at the first 4 hours, signifying the presence of NiO and the production of BnNH2 was hindered.

However, using the activated catalyst, very little PhCN can be seen throughout the reaction, and the catalyst can reach full potential to producing primary amines selectively.

It is interesting that PhCN was not observed in previous batch reactions, despite the same unreduced catalyst being used. This is because the effective catalyst to reactant ratio in the flow system is much higher than in batch, thus quickly converting the reactants to the irreversible PhCN.

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6.4. Conditions Study

Reaction conditions can be more easily controlled in a flow setup compared to batch, as discussed in Chapter 1.5 (Flow Reactors). In this study, a number of conditions were investigated. These conditions include:

• Reaction temperature,

• Catalyst particle size,

• NH3/BnOH molar ratio,

• Flow rates, and

• Initial alcohol concentrations

For each condition, a set of reactions of NH3 N-alkylation with BnOH were carried out and results were analysed in terms of the % conversions of BnOH and selectivities to

BnNH2. The pressure was kept constant at 60 bar for all reactions.

6.4.1. Reaction temperature

Increasing the reaction temperature should increase reaction rates. However, this increase in reaction rate may affect the selectivity of the amine products because they have different energies of formation.

Reactions with temperatures from 120 to 200 °C were carried out with the flow reactor with 0.6 M BnOH solution and NH3/BnOH = 7 at a flow rate of 0.5 mL/min, which accounts to a residence time of 5.6 min. The results are shown in Table 6.1.

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Table 6.1: Flow reaction results from 120 to 200 °C.

Temperature

Entry (°C) Conv. (%) Sel. to BnNH2 (%)

1 120 9% 100%

2 160 38% 100%

3 200 92% 91%

Conditions: 0.5 mL/min, NH3/BnOH = 7, 2.0 g Catalyst 2, residence time = 5.6 min, 60 bar, T = 120 – 200 °C, Conv. = Conversion; Sel. = Selectivity.

The conversion was low at 120 °C with 9%, but a high selectivity to BnNH2 was achieved at 100%. As the reaction temperature was increased to 160 °C, the conversion increased to 38% and the selectivity to BnNH2 remained high at 100%. For 200 °C, a high conversion of 92% was achieved, but the selectivity to BnNH2 decreased to 91% with a selectivity of 8% towards Bn2NH.

It was observed that the higher the temperature, the higher the conversion, although causing slightly lower selectivity to BnNH2. The increase in conversion is expected because a higher temperature promotes a higher reaction rate. This is because of a higher rate of collisions between molecules, and more molecules would possess the energy required to overcome the activation energy to form BnNH2 or Bn2NH. As a result, the increase in temperature to 200 °C may have promoted the formation of Bn2NH, thus decreasing the selectivity to BnNH2.

160 °C was chosen as the operating temperature for flow reactions. This is because the high selectivity to BnNH2 is highly advantageous for product separation. The low conversion of 38% can be increased by increasing the residence time of the reaction solution, which is investigated further in Chapter 6, Section 6.5.4.

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6.4.2. Catalyst size

The smaller catalyst size has a higher surface area, which allows more active sites to be exposed to the reaction solution, thus increasing conversion. It is unclear if the selectivities would change because the nature of these active sites may change after the grinding stage.

Smaller catalysts (250 – 600 µm, Catalyst 3) were produced by grinding down the 2 mm extrudites (Catalyst 2), then sieved and washed with de-ionised (DI) water to remove fragments. 2.0 g of catalyst was employed for the flow reactions, and the results are shown in Table 6.2. The catalyst bed length should be 0.25 m, but the smaller size of

Catalyst 3 caused a tighter packing. Therefore, 0.5 g of inert bentonite extrudites were added to Catalyst 3 and dry-mixed before the packing to ensure an equal distribution of the catalyst bed.

Table 6.2: Flow reaction results at different catalyst grades.

Catalyst size Sel. to amines (%)

Entry Name (mm) Conv. (%) BnNH2 Bn2NH

1 Catalyst 2 2 38 100 0

2 Catalyst 3 0.25 – 0.60 20 84 16

Conditions: 0.5 mL/min, NH3/BnOH = 7, 2.0 g catalyst, residence time = 5.6 min, 160

°C, 60 bar, Conv. = Conversion; Sel. = Selectivity.

Surprisingly, the conversion was lower at 20% for Catalyst 3 compared to 38% for

Catalyst 2. There was also a lower selectivity to BnNH2 of 84%, and Bn2NH as the only side products at 16%. This may suggest that the crushing of the catalyst to finer particles may lead to the mechanical failure of active sites, as well as the change in selectivity.

A review done by Trimm et al. [161] regarding the control of pore size in alumina catalyst supports suggested many catalyst deactivation mechanisms, such as poisoning,

174 fouling and thermal degradation. However, these types of deactivation mechanism should be seen with both catalysts. Another mechanism that may lead to deactivation is the different forms of mechanical catalyst failure which involves the following factors:

1. Crushing of granular, pellet or monolithic catalyst forms;

2. Attrition, the size reduction and/or breakup of catalyst granules and pellets to

produce fines, particularly in fluid or slurry beds; or

3. Erosion of catalyst particles or monolith coatings at high fluid velocities.

The crushing of catalyst pellets to smaller particles were carried out prior to catalyst packing and reaction, so 1 is a possibility. 2 is unlikely because there should be no movement allowed in the packed bed reactor. 3 is also unlikely because the flow rate of the reaction solution was not high at 0.5 mL/min. Therefore, the mechanical crushing may have caused the decrease in conversions. To confirm this hypothesis, Brunauner-

Emmett-Teller (BET) specific surface area measurements was conducted on both catalysts prior to reactions and the results are shown in Table 6.3.

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Table 6.3: Brunaun-Emmett-Teller (BET) specific surface area measurements of Ni

catalysts of different sizes.

BET

Catalyst size Surface Area (SBET) Average pore Pore size

2 3 Entry Name (mm) (m /g) volume (cm /g) (dp.Avg) (nm)

1 Catalyst 2 2 127 0.371 11.7

2 Catalyst 3 0.25 – 0.60 34.6 0.064 7.44

Catalyst 2 had a surface area of 127 m2/g, with an average pore volume of 0.371 cm3/g

and a pore size of 11.7 nm. Catalyst 3 had a surface area of 34.6 m2/g, with an average

pore volume of 0.064 cm3/g and a pore size of 7.44 nm. There was a decrease in surface

area, pore volume as well as the average pore diameter. The surface area had decrease

by almost 4 times, the pore volume decreased by 6 times and pore size by 36%. This

could confirm that the crushing of catalyst caused the mechanical distortion of the

catalyst.

As there was no advantage in using a finer particle size catalyst, the subsequent reactions

were performed using 2 mm catalyst pellets (Catalyst 2).

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6.4.3. Ammonia/Alcohol Molar Ratio

It is expected that increasing the ratio of NH3/BnOH increases the selectivity to primary amines. This is because when considering the following reactions, they are competing as BnOH is the starting material for both reactions, with Reaction 2 being the faster reaction as discussed in Chapter 5, Section 5.4 (Kinetic Studies of Batch Reactions):

[BnOH][NH3] ⇆ [BnNH2][H2O] (Reaction 1)

[BnNH2][BnOH] ⇆ [Bn2NH][H2O] (Reaction 2)

An increase in [NH3] encourages a faster reaction 1 than reaction 2, thus increasing the selectivity to BnNH2.

Experiments with different ratios of NH3/BnOH were carried out from 3 to 10 to achieve the highest selectivity to BnNH2. The mixture had a combined flow rate of 0.5 mL/min over different NH3/BnOH ratios. Therefore, individual flow rates of the components were calculated in advance as shown in Chapter 3, COFE simulations, which accounted for the density of liquid NH3. It should be noted that in a flow setup, the stoichiometry can be easily adjusted by simply changing individual flow rates of the reaction components, which is advantageous compared with a batch setup. The reaction conditions were a combined reactant flow of 0.5 mL/min, 60 bar, 160 °C, 0.6 M BnOH initial concentration, 2.0 g Catalyst 2 and NH3/BnOH = 3, 5, 7, 10. Results are shown in Table 6.4.

As the NH3/BnOH molar ratios increased from 3 to 7, conversions increased linearly from 14% to 41%, but only had a slight increase to 44% as NH3/BnOH increased to 10.

This observation suggests that the reaction rate is proportional to [NH3] up to

NH3/BnOH = 7. Beyond that, the conversion becomes independent of [NH3], as NH3 is in excess in the system and [NH3] is considered a constant.

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Table 6.4. Flow reaction results at different NH3/BnOH ratios.

Entry NH3/BnOH Molar Ratio Conv. (%) Sel. to amines (%)

BnNH2 Bn2NH Bn3N

1 3 14 0 84 16

2 5 27 79 21 0

3 7 41 67 32 1

4 10 44 77 23 0

Conditions: 0.5 mL/min, 2.0 g Catalyst 2, residence time = 5.6 min, 160 °C, 60 bar,

NH3/BnOH = 3 – 10, Conv. = Conversion; Sel. = Selectivity.

The selectivities of the products changed dramatically from NH3/BnOH = 3 to 5 but remained similar for 5 to 10. At NH3/BnOH = 3, there were no BnNH2, but 84% of

Bn2NH and 16% of Bn3N; whereas for NH3/BnOH = 5 – 10, the reactions are more selective towards BnNH2 at 67 – 79% with Bn2NH as the side product, with very little production of Bn3N.

In the case of NH3/BnOH = 3, the high selectivity of the Bn2NH can be attributed to

Reaction 2 happening at a faster rate than reaction 1, which is already known from the reaction kinetics described in the previous chapter. For NH3/BnOH = 5 – 10, the selectivity for BnNH2 is much higher because [NH3] was high.

It was found that selectivity and conversions increase as NH3/BnOH increases from 3 to 7, but did not change much from NH3/BnOH = 7 to 10. The following reactions were carried out at NH3/BnOH = 7 because the increase at 10 was incremental. Additionally, it was safer to operate at NH3/BnOH = 7, as it would require a higher pressure to operate at the liquid phase at NH3/BnOH = 10, as demonstrated in previous COFE simulations.

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6.4.4. Flow Rates

It is expected that as residence time increases, the conversion of BnOH would increase, but selectivity to BnNH2 would decrease as more Bn2NH would form. This is because

BnNH2 is more nucleophilic than NH3, and would compete to react with the aldehyde, as discussed in Chapter 5. The reaction time in flow, i.e. the flow rate must be carefully selected to achieve high selectivity to BnNH2.

Reactions of flow rates = 0.1 – 1 mL/min were performed and the conversions and selectivities measured. The void volume of the packed bed was 2.8 mL, so the residence times were calculated as 2.8 – 28 min. The initial concentration of BnOH solution was

0.6 M in o-xylene. The reaction conditions were 160 °C, 60 bar and 2.0 g of Catalyst 2.

The conversion and selectivity to BnNH2 are shown in Figure 6.6.

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The conversion was modest at 58% at 0.1 mL/min, which increased slightly to 51% at

0.2 mL/min, but dropped to 12% at 1 mL/min. This is expected because when the flow rate is faster, residence time is lower, which means a shorter reaction time.

Selectivity to BnNH2 increased from 30% at 0.1 mL/min to 85% at 1 mL/min. This is because the faster flow rate has allowed BnNH2 to pass through the reactor more quickly without further reacting to form Bn2NH.

As expected, when the flow rate was increased, the conversion decreased, but the selectivity to BnNH2 increased. It is difficult to strike a balance for such a high concentration as the outcomes were not satisfactory to achieve high conversions.

Therefore, lower BnOH concentrations were explored at 0.1 and 0.2 M.

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6.4.5. Initial Alcohol Concentrations

It is expected that the lower the initial concentration, the lower the rate of reaction.

However, the increase of Ni/BnOH ratio should increase the BnOH conversion.

A lower concentration of BnOH solution at 0.2 M in o-xylene were tested in the flow reactor under conditions of 60 bar, 160 °C, 2.0 g Catalyst 2, with flow rates of 0.1 to 1 mL/min. Results are shown in Figure. 6.7.

Figure 6.7. Plot of conversion of 0.2 M BnOH and selectivity to BnNH2 under different flow rates. Conditions: 0.1 – 1 mL/min, 2.0 g catalyst, residence time = 2.8 - 28 min,

160 °C, 60 bar 0.2 M BnOH in o-xylene, NH3/BnOH = 7, 2.0 g Catalyst 2.

For reactions with initial concentration of BnOH = 0.2 M, conversion was high for a low flow rate of 0.1 mL/min at 78%, but dropped with higher flow rates to 32% at 1 mL/min, due to the lower residence time. On the other hand, the selectivity to BnNH2 remained high in the range of 86 – 100%.

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The concentration of BnOH solution was further lowered to 0.1 M and the reaction was carried out in the flow reactor under conditions of 60 bar, 160 °C, 2.0 g catalyst B, over flow rates of 0.06 – 0.5 mL/min. Results are shown in Figure. 6.8.

Figure 6.8. Plot of conversion of 0.1 M BnOH and selectivity to BnNH2 under different flow rates. Conditions: 0.06 – 0.5 mL/min, 2.0 g catalyst, residence time = 2.8 – 47 min,

160 °C, 60 bar 0.1 M BnOH in o-xylene, NH3/BnOH = 7, 2.0 g Catalyst 2.

For reactions with an initial concentration of BnOH = 0.1 M, conversions were high at lower flow rates of 0.06 mL/min at 90%, which dropped linearly at higher flow rates of

0.50 mL/min to 22%. This is expected as conversion should be directly proportional to the residence time. Selectivity was consistently high at a remarkable 100% over the flow rates of 0.06 – 0.5 mL/min, which is more uniform than higher concentrations.

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Figure 6.9. Plot of conversion of 0.1 M, 0.2 M and 0.6 M BnOH to BnNH2 under different flow rates. Conditions: 0.06 – 1 mL/min, 2.0 g catalyst, residence time = 2.8 –

47 min, 160 °C, 60 bar 0.1 M BnOH in o-xylene, NH3/BnOH = 7, 2.0 g Catalyst 2.

By comparing results of 0.1 – 0.6 M, it is apparent that the lower the flow rate, the higher the conversion was achieved. The selectivity to BnNH2 also increased as the concentration decreased. More specifically, 100% selectivity could be achieved with a low concentration of 0.1 M. A high selectivity can aid downstream processes with fewer workup. Thereafter, BnOH solutions were prepared at 0.1 M for the flow reactions, unless stated otherwise.

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6.4.6. Summary for Reaction Conditions

The following conditions were investigated for the N-alkylation of NH3 with BnOH using a flow reactor.

• reaction temperature = 160 ºC,

• reactant molar ratio at NH3/BnOH = 7,

• catalyst size = 2 mm,

• flow rate = 0.06 mL/min, and

• initial alcohol concentration = 0.1 M

There is a fine balance for all the conditions in order to achieve the desired outcome of the reaction between NH3 and BnOH, with a goal to achieve high selectivity to BnNH2.

One advantage with the flow setup is that it allowed a quick and easy way in adjusting these conditions for improvements.

When compared to batch reactions, the reaction temperature was the same at 160 ºC, but the flow reactor enabled a higher reaction temperature at 200 ºC under safe conditions, but the higher temperature was shown to lower the selectivity to primary amine, which is undesirable. The NH3/BnOH molar ratio was 7, which was the same as batch conditions. The flow reactor has enabled a higher NH3 concentration under safe conditions, but was found that NH3/BnOH molar ratios higher than 7 does not increase the selectivity. The catalyst size in the flow reactor was larger than that of batch reactors

(2 mm compared to 250 µm, and the reaction times in flow were shorter than that of batch reactors (47 minutes compared to 72 hours). Lastly, the initial alcohol concentration of flow reactions is lower than batch reactions (0.1 M compared to 0.6

M), as it was found that the higher concentrations promoted formation of Bn2NH.

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The following reaction conditions have been chosen to prioritise selectivity to BnNH2, which is 160 °C, 2.0 g Catalyst 2, NH3/BnOH = 7, 0.06 mL/min flow rate and initial concentration of 0.1 M. These reaction conditions will be used in the following experiments, unless stated otherwise.

6.5. Flow Reactions with Different Alcohols

Flow reactions with different alcohols were carried out with the conditions stated above.

The results are shown in Table 6.5. The results from the flow reactions are then compared with the batch reactions.

Table 6.5: Fixed bed flow reactions results with various alcohols.a Entry Starting material Product Conv. Sel. to primary

(%) amine (%)

1 100 99

2 86 99

3 100 99

4 51 99

5 67 99

6d 89 94

7d 78 90

a Conditions: 0.10 M ROH in o-xylene, NH3/ROH = 7, 2.0 g extruded 36 wt% Ni-

b Al2O3/SiO2, 0.06 mL/min, residence time (τ) = 46.7 min; Steady-state achieved after

15 x τ and maintained for 3 x τ. c0.03 mL/min, τ = 93 min. dKetone detected as byproducts. eSecondary amine detected as byproducts.

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Using a BnOH solution with a concentration of 0.10 M, flow rate of 0.06 mL/min, 160

°C and NH3/BnOH stoichiometry of 7, 100% conversion and selectivity for BnOH could be achieved, (Table 6.5, Entry 1) compared to the conversion of 96% and selectivity of

97% in a batch system (Chapter 5). These flow reaction conditions were carried forward for the other alcohols with significant improvements in conversions and selectivities compared to the batch reactions.

The methoxybenzyl alcohols had higher conversions in the flow system compared with the batch system. (Table 6.5, entries 2, 3) 2-methoxybenzyl had a conversion of 86% in flow compared to 26% in batch, and 3-methoxybenzyl had a conversion of 100% in flow compared to 24% in batch. The reason why 2-methoxybenzyl alcohol achieved a lower conversion might be steric hindrance from the 2-methoxy group, despite the electron donating effect at the ortho- position. 3-methoxybenzyl matched the performance of benzyl alcohol because the meta- position is not affected. The selectivities to primary amines for both 2- and 3-methoxybenzyl alcohol were 99%, which is the same as the batch results.

Scheme 6.3. Steric effects of a 2-methoxybenzaldehyde (left) compared to a 3-methoxybenzaldehyde during the condensation reaction with NH3.

High conversions were achieved with picolyl alcohols (Table 6.5, entries 4 & 5) when compared to the batch reactions. 2-picolyl alcohol had a conversion of 51% in flow compared to 18% in batch, and 3-picolyl alcohol had a conversion of 67% in flow compared to 41% in batch. The increase in conversions was not as high as that of the methoxybenzyl group, possibly due to leaching and thus lowered catalytic activity. This is because picolyl alcohols are more polar than benzyl alcohols and may induce leaching

186 nickel, especially with 2-picolyl alcohol as the two nitrogen atoms may form ligands with Ni. However, the eluent was analysed by ICP and the Ni content was under the detection limit (>20 ppm). Selectivities to the primary amines remained high at 99%.

The alkyl alcohols experiments were performed under a lower flow rate of 0.03 mL/min, which was half the flow rate for aryl alcohols (Table 6.5, entries 6, 7). 2-octanol achieved a conversion of 89% compared to 10% in batch conditions, and 1- phenylbenzylalcohol achieved 78% conversion compared to 7.8% in batch. The selectivities to the primary amines were higher than the batch reactions at 94% and 90% respectively, compared to 0% and 55% of batch conditions. The difference between the batch and flow results is attributed to the higher effective concentration of catalyst under flow conditions, thus the alkyl alcohols performed better in flow reactors.

Shimizu et al. [122] achieved a Turnover Number (TON) at 15.6 for 72 hours in batch when using benzyl alcohol. Although a lower TON of 0.25 over 27 hours was achieved in this study, this system has the advantage of being in flow and is a continuous production. In addition, this system was optimised to obtain 99% selectivity for BnNH2, where Shimizu et al. [122] achieved a lower selectivity of 78%. A high selectivity could be more advantageous as it can potentially save costs in product separation and purification.

6.6. Comparisons with Different Primary Amine Production Methods

As discussed in Chapter 2, Sections 2.3 – 2.8 (Literature Review), there are currently six methods in primary amine production. However, most of these processes are carried out in batch reactions and are thus difficult to compare. The following example of nitrile hydrogenation was carried out in flow with supported Pd catalyst as reported by

Kobayashi et al. [102]. A low hydrogen pressure of 0.5 bar was used in their flow reactor, which was heated at mild temperature of 60 °C. High yields of the ammonium salt were achieved at 97 – 100%. The catalyst remained active for more than 300 hours

187

(TON > 10 000 mmol ammonium salt/mmol Pd) without loss of selectivity. The TON of this Pd catalyst is much higher than Catalyst 2 (0.25 mmol BnNH2/mmol Ni over 27 hours). The use of Pd catalyst is more expensive and the separation of ammonium salt would require additional workup downstream.

Scheme 6.4. Flow scheme of selective hydrogenation of nitriles to primary amines catalysed by a supported Pd catalyst, reported by Saito et al. [102].

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6.7. Extended flow reaction over time

In a prolonged flow experiment over 80 hours, 0.10 M BnOH solution was reacted with

NH3 (7 eq.) at 160 °C under 60 bar over Catalyst 2 at a steady 0.06 mL/min flow rate.

The TON and selectivity to BnNH2 over time is shown in Figure 6.10.

BnNH2 was produced from the beginning with 100% selectivity except for the first 5 hours.

The TON (in square symbols, Figure 6.10) was calculated with regards to the accumulated mmol of benzyl amine produced per mmol basic sites. It shows a slow increase at the beginning, accelerates from 30 hours and becomes linear at 60 hours, but slows down and reaches, reaching 7800 mmol amine/mmol basic sites at 80 hours. This might suggest that the catalyst is decreasing in activity after 60 hours.

Brunauer-Emmett-Teller (BET) and Temperature-Programmed Desorption (TPD) analysis were carried out on Catalyst 1, bentonite, Catalyst 2 before and after the reaction, and the results are shown Table 6.6.

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Table 6.6. Ni catalysts at different stages and their BET & TPD results.

BET TPD

Average

SBET pore volume dp.Avg Basicity Acidity

Sample (m2/g) (cm3/g) (nm) (µmol/m2) (µmol/m2)

Catalyst 1 53.1 0.184 17.0 13.1 22.5

Bentonite 90.3 0.352 13.5 5.8 nil

Catalyst 2 127 0.366 11.0 12.9 1.8

Catalyst 2 after reaction 152 0.412 10.2 23.7 3.6

BET: Total surface area of Catalyst 2 increased after combining the Catalyst 1 with bentonite and increased further post-reaction. The pore volume of Catalyst 2 was much higher than Catalyst 1 but comparable to bentonite, which increased after reaction.

However, the pore diameter of Catalyst 2 is lower than Catalyst 1 and bentonite, which decreased slightly after reaction.

The differences between Catalyst 2 before and after reaction could be due to meso-pores expansion during the reaction as the substrates can cause changes to the catalyst structure when chemisorbed on the catalyst surface.

TPD: Basic site density of Catalyst 2 was comparable with original catalyst at 12.9 ±

0.9 µmol/m2 vs 13.1 ± 3.4 µmol/m2. However, the density of acidic sites decreased a lot.

This might be due to the lack of acidic sites in bentonite. It is interesting to see acidity/basicity almost doubled after reactions. This may be due to more defective sites produced as substrates were passed through the catalyst.

These results are not conclusive to explain the decrease in catalyst activity over time, and thus a different technique is used in the following section.

190

6.8. Catalyst deactivation

The long-term stability of a catalyst is critical to its industrial viability, as costs of replacements could out-value the low market prices of base metals. Catalyst activity degrades due to the loss of active sites, either by the loss of nickel content (leaching), or the active sites being blocked (poisoning).

As demonstrated in Figure 6.10, the steady-state conversion of benzyl alcohol to benzyl amine decreased slowly over a period of time (after 72 h), which suggests catalyst deactivation. In this part of the work, the spent catalysts from flow or batch reactions were examined using ICP and TGA.

The absence of Ni content in all liquid phases (ICP analysis <0.5 mg/L), suggested that catalyst leaching is not a significant issue in the absence of a Lewis basic group (such as a pyridyl group in the picolyl substrates, as described in Chapter 4). Instead, TGA-

MS experiments performed with the recovered catalysts showed volatile compounds in the form of carbon dioxide (m/z = 44) emerging at 300 °C, indicating that the catalyst deactivation was a result of carbonaceous deposition of organic compounds onto the catalyst surface. The C:H ratio of the gaseous thermal product was 1:1.2, suggesting a form of aromatic oligomer has been formed, possibly catalysed by bentonite, the co- support used in the extrusion method. As Morales-Serna et al. [162] suggested, a montmorillonite clay is able to catalyse the oligomerisation of 3,5-dimethyl benzyl alcohol. However, this analysis is inconclusive and would require future work to confirm the proposed deactivation pathway. Details are discussed in Chapter 7, Section

7.2, Recommendations for Future work.

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6.9. Conclusion

In this chapter, N-alkylation of NH3 with alcohols was performed using a continuous flow system.

Firstly, a bench-top continuous flow system was re-designed and constructed for the N- alkylation of NH3 with alcohols. The system could withstand high pressure up to 80 bar to contain NH3 as a liquid. The fine commercial catalyst (65 wt% Ni/Al2O3/SiO2) was extruded prior to administration in the plug flow reactor to reduce pressure drop, and an in-line filter was fitted. The reactor system was characterised and it was confirmed that the reaction mixture was under liquid phase during the flow reaction using two simulation methods.

During a preliminary flow experiment an unexpected product (benzonitrile) was observed. That was due to the presence of nickel oxide (NiO) in the catalyst extrudites, which was resolved by a catalyst prereduction method.

After investigating several reaction conditions, including reaction temperature, catalyst size, NH3/BnOH molar ratios, flow rate and initial alcohol concentration, the conditions which had the highest BnOH conversion and BnNH2 selectivity were as follows: 160°C,

2.0 g 2-mm extrudite catalyst, NH3/BnOH = 7, flow rate = 0.06 mL/min and initial alcohol concentration of 0.1 M, which achieved 100% BnOH conversion and 99%

BnNH2 selectivity. N-alkylation of NH3 with seven other alcohols were achieved in the flow system, with 90 – 99% conversions and 51 – 100% selectivities to the respective primary amines. This was a significant improvement compared to batch reactions, where conversions were much lower at 18 – 41%.

The system was able to operate for 80 hours at a selectivity to BnNH2 at 99%, yielding a TON of 7800 mmol amine/mmol basic sites on catalyst. Catalyst deactivation over long periods of reaction time was observed and was attributed to carbonaceous deposition.

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Chapter 7: Summary and Recommendations for

Future Work

7.1. Summary

The aim of this project was to develop a new process to produce primary amines with the following aspects: (i) hydrogen borrowing cycle, (ii) safe working environment, and

(iii) atom efficiency. The experimental study addressed the following objectives: (i) understanding the thermodynamics of the system, (ii) finding a suitable catalyst, (iii) studying reaction kinetics to control rate and selectivity, and (iv) conceiving, designing and constructing a continuous flow system that allows safe handling of ammonia.

Overall, a highly selective N-alkylation of ammonia with alcohols to primary amines was achieved using a commercially available catalyst (65 wt% Ni/Al2O3/SiO2), both in batch and flow reactors.

(i) Understanding the thermodynamics of the system,

The thermodynamics of the reaction system was investigated in two parts. Firstly, the reaction outcome under different reactant stoichiometry was investigated using ASPEN- plus. The reaction outcome of different BnOH:NH3:H2O ratios were calculated under different temperatures and pressures. It was found the temperature and pressure had very little effect on the reaction outcome, and that upon reaction completion, BnOH conversion always reaches 100%. The product selectivity towards the primary amine and tertiary amine was changed by increasing NH3 or BnOH respectively, and H2O affected very little of the reaction outcome, despite being a by-product in the condensation step of the hydrogen borrowing technique.

Secondly, the physical state of the reaction mixture under the flow system was calculated. Two simulation methods (COFE and ASPEN-plus) confirmed that the

193 reaction mixture was under liquid phase under the reaction conditions of 160 °C and 60 bar.

(ii) Finding a suitable catalyst,

Using the N-alkylation of NH3 with BnOH as a model reaction in the batch system, different commercial catalysts were first screened, including Ru, Pt, Au and Ni. Au and

Pt had not been previously demonstrated to catalyse this reaction. It was found that the

Ni catalysts were selective towards the primary amine at 85 – 96%, with 65 wt%

Ni/Al2O3/SiO2 having the highest BnOH conversion at 44%. On the other hand, the noble metals, Ru, Pt and Au were selective towards the higher amines, specifically with

1 wt% Au/TiO2 having a high BnOH conversion at 48% and high selectivity towards the secondary amine at 60%.

Ni catalysts with different Ni loadings were synthesised using a coprecipitation method.

Although the synthesised catalysts had high selectivity to BnNH2 at 75 – 94%, the conversions were much lower than the commercial catalyst at 18 – 32% compared to

96%. It was found that NiO was the predominant species on the synthesised catalyst, even after reduction under H2, while Ni is the predominant species on the commercial catalyst, hence the difference in activities.

The extrusion method was applied to the 65 wt% Ni/Al2O3/SiO2 catalyst and has enabled the application of the commercial catalyst in the continuous flow system.

(iii) Studying reaction kinetics to control rate and selectivity,

The batch reaction mixture was monitored over time and the reaction rate constants were determined using the experimental data by Berkeley Madonna. Au had higher kinetic rate constants for the N-alkylation of the primary and secondary amines with alcohol than Ni, which explains the higher selectivity to secondary and tertiary amines.

194

Therefore, Ni was selected as the main catalyst, as it was highly selective towards the primary amine.

Batch reactions with intermediate amines and water were performed with the goal of improving the reaction model. A three-reaction and a four-reaction kinetic model were developed and compared but was found to be equally suitable.

Batch reactions with different NH3 and H2O stoichiometries were performed and it was proven that water has an adverse effect on the forward reaction. Thus, an NH3 condensation method was developed for the introduction of anhydrous NH3 into batch reactors, which has greatly improved the conversions from 27% to 96%, while the selectivity to BnNH2 remained high at 97%. By plotting the Arrhenius graph, the activation energy, Ea and the frequency factor, A of the monoalkylation of NH3 with

BnOH using the Ni extrudites were found to be 98 kJmol-1 and 9.7e11 s-1, respectively.

(iv) Conceiving, designing and constructing a continuous flow system that

allows safe handling of ammonia and achieving higher yields than the

batch system

A bench-top continuous flow reactor was re-designed and constructed for the N- alkylation of NH3 with alcohols, where NH3 could be safely delivered as a reagent.

Specifically, an NH3 condensation unit was installed at the beginning for the supply of anhydrous NH3 as a liquid, a gas-liquid separation unit and an acid trap was installed at the outlet. The fine commercial catalyst (65 wt% Ni/Al2O3/SiO2) was extruded prior to administration in the plug flow reactor, and an in-line filter was fitted.

An unexpected product (benzonitrile) was observed because NiO was in the catalyst extrudites, which was resolved by a catalyst prereduction method. The conditions were:

160 °C, 2.0 g of 2-mm extrudite catalyst, NH3/BnOH = 7, flow rate = 0.06 mL/min and initial alcohol concentration of 0.1 M, which achieved 100% BnOH conversion and

99% BnNH2 selectivity. The system was able to operate for 80 hours at a 99% selectivity

195 to BnNH2, yielding a TON of 7800 [mmol amine]/[mmol basic sites on catalyst].

Catalyst deactivation over long periods of reaction time was observed and was attributed to carbonaceous deposition.

In addition to the stated objectives, the scope of alcohol substrates was expanded to see how different alcohols behave under the system. This included other aromatic alcohols such as picolyl alcohols, as well as alkyl alcohols such as 2-octanol and 1,4-butandiol.

These experiments were performed in batch and flow reactors using the same catalyst extrudites. While most aromatic alcohol substrates performed moderately in the batch system with low conversions at 18 – 41% and high selectivities to their corresponding primary amines at 96 – 99%, the performances of alkyl alcohols in batch were inferior at conversions of 8 – 10% and 0 – 55% selectivity to primary amines. On the other hand, these alcohols performed well in the flow system, with alcohol conversions at 51 – 99% and selectivities to their corresponding primary amines at 90 – 99%.

In the following section, certain aspects of the project are suggested for any future work, as some objectives were not fully realised, or improvements could be made to the current setup.

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7.2. Recommendations for Future Work

1. Complete batch reactions

According to the thermodynamics calculations in Chapter 5.2 (Thermodynamics), the reaction should reach 100% completion with 100% conversion of BnOH in most cases.

However, the experimental reaction times were insufficient to achieve complete reactions. Further work may include longer reaction times or higher catalyst loading to drive the reaction to completion and compare with the current work.

2. Solvent system o-xylene was used extensively in the project as a solvent due to its high boiling point and good solubility of the model alcohol, BnOH. However, the more polar alcohols such as the diols used in Chapter 5.8 (Batch System with Different Alcohols) were unable to fully dissolve in o-xylene, and thus were not used in the flow system. Developing a more polar solvent system such as using tetrahydrofuran can allow the dissolution of more polar alcohols such as diols or even glycerol, a renewable platform chemical [163].

However, this may cause issues in catalyst leaching. Therefore, the development of catalysts that can withstand polar conditions would be beneficial.

3. Ammonia source

The use of anhydrous ammonia instead of aqueous ammonia in the batch system was crucial in achieving high conversions, as it was proven that water inhibits the forward reaction in chapter 5.6 (Effect of Water). Alternatively, non-volatile ammonia substitutes such as urea, ammonium nitrate or ammonium formate can be used in place of anhydrous ammonia, as demonstrated by Yamaguchi et al. [121] in the N-alkylation of ammonia with BnOH using a Ru catalyst. It was found that urea gave the highest yield to the tertiary amine at 93%, followed by aqueous ammonia at 87%. The advantage of using a non-volatile ammonia substitute is that the system need not be pressurised.

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However, when compared to anhydrous ammonia, these substitutes are less atom efficient as they would generate waste. These substrates are also more difficult to recycle, and anhydrous ammonia can be easily separated from the reaction stream as a gas at 1 bar.

4. Kinetic studies for other catalysts and application in flow

Ni and Au were selected for kinetic studies due to high selectivities to BnNH2 and

Bn2NH respectively. It was proven in Chapter 5.4 (Kinetic Studies) that the reaction rates of Au is much faster than Ni, thus the higher selectivity to Bn2NH and Bn3N. By controlling the reaction conditions, such as increasing the flow rate or NH3 concentration, a highly selective synthesis of primary amine may be achieved using Au in flow.

It would be interesting to expand the scope of catalysts and investigate the kinetic constants for other catalysts such at Ru and Pt, where high selectivity to Bn3N were achieved in the preliminary experiments at 63% and 36% respectively (Chapter 5.3.

Preliminary Batch Reactions). These catalysts may also be applied in the flow reactor, where the product selectivity may be controlled by manipulating the reaction conditions.

5. Catalyst extrudite size control

As previously demonstrated in Chapter 6.4.2 (Flow Reactions, Condition studies,

Catalyst size), the catalyst size of the catalyst extrudites were reduced by grinding, which caused pore collapse and deformation of active sites, as shown in the BET results.

To prevent this, catalyst with a higher attrition resistance and strength can be synthesised, which can undergo harsher treatment without collapsing, thus allowing access to smaller catalyst sizes and increasing the surface area. Aggregate/agglomerate strength can be increased by means of advanced preparation methods, such as sol-gel granulation, spray drying with organic binder, or controlled precipitation methods [164].

198

6. Analysis methods

A different analysis approach may be used for batch reactions. The current GC-FID analysis required sampling of the reaction mixture over time, which decreased the reaction volume. The use of Nuclear Magnetic Spectroscopy (NMR) analysis may resolve this, where the alcohol and amine peaks can be tracked over time. However, this is difficult to achieve using the current experimental setup with anhydrous ammonia being toxic and volatile. Thus the use of a non-volatile ammonia substitute would aid this endeavour.

Theoretically, the flow reaction of N-alkylation of NH3 with alcohols can be monitored online using infra-red (IR) detection, as the alcohol O-H stretch and amine N-H stretch(es) can be detected in the absorption of 3550 – 3200 cm-1. However, the alcohol and amine peaks overlap, making the determination of individual species difficult.

Realistically, the IR detector can only track the consumption of BnOH over time as an online analytical tool, as the O-H absorption is much stronger than N-H.

An inline titration method can be used in the flow reactor to verify the actual concentration of NH3. An inline titration method developed by Pastre et al. [165] used an injection of a dilute acid containing a pH indicator to the reaction mixture and was used to measure the effect of temperature on NH3 uptake in different solvents at -20 –

80 °C at 6 bar. This analytical method can be replicated for the current flow system, where harsher conditions were applied.

An online gas phase analysis of the outlet for both batch and flow reactions would provide insights on the reaction mechanism, specifically for H2 and carbon containing molecules. If H2 gas was not detected, this would help confirm the first step in the mechanism of the hydrogen borrowing cycle, as well as confirm the formation pathway of the secondary imine that was detected with reactions involving amine intermediates and water (Chapter 5, Section 5.5). Also, the discovery of any carbon containing gas

199 molecules may shed light to any unexpected reaction pathways, which would deepen the understanding of the overall reaction mechanism.

As mentioned in Chapter 6, Section 6.8, Catalyst Deactivation, the analytical results from the TGA-MS might not be conclusive enough to indicate oligomerisation of benzyl alcohol, catalysed by the montmorillonite clay. NMR analysis might provide insight on to the unknown carbonaceous deposition on the catalyst. Furthermore, a revival of the deactivated catalyst can be achieved by heating the catalyst under air, followed by H2.

7. Scaling up of flow system

An advantage of continuous flow system over batch system is the ease to scale up.

Scaling up in volume of a batch reactor has implications on the heat and mass transfer within the system, and the reaction conditions must be re-optimised for the larger scale.

For example, Styring and Parracho [166] used a parallel capillary reactor for the scale up of a cross-coupling meso-flow reaction from discovery phase to production phase of chemical synthesis. By replication of the single reactor in parallel, they achieved the same chemistry on a larger scale, on a small footprint and without the heat and mass transport limitations of batch reaction scale-out.

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Appendix

Properties of ammonia

Ammonia phase diagram

Figure S1. Phase diagram of ammonia, from pressures between 0 to 50 bar. The plots are vapour pressures of the gas at different temperatures Lange's Handbook of

Chemistry, 10th ed. page 1451 and 1468.

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Ammonia Solubility in Toluene

Ammonia behaves differently from an ideal gas, as the vapour pressure and temperature do not have a linear relationship. The vapour pressure is 10 bar at 25 °C.

Figure S2. NH3 solubility in toluene at 120 °C, Taken from Bhattacharyya et al. [167].

Scheme S1. Synthetic route of Lisinopril [3].

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S3. Thermodynamic outcome of BnOH/NH3/H2O from ASPEN presented in mol%

The reaction conditions were first set at a pressure of 10 bar and temperature of 160 °C

(Table S1, Entries 1 – 5), then at pressure 60 bar and temperatures of 160, 120 & 80 °C

(Table S1, Entries 6 – 10, 11 – 15 & 16 – 20 respectively). These conditions were tested

with different BnOH:NH3:H2O ratios: 1:1:1, 1:1:10, 1:1:100, 1:1:0, 6:1:0, 1:3:0.

Table S1. A prediction of the thermodynamic outcome from ASPEN, presented in mol%. Conditions Input (mol%)

Entry Output (mol%)

Temperature (°C) Pressure (bar) BnOH NH3 H2O BnNH2 Bn2NH Bn3N

33% 33% 33% 0% 0% 0% 1 0% 4% 66% 27% 1% 2%

8% 8% 83% 0% 0% 0% 2 0% 1% 92% 7% 0% 0%

1% 1% 98% 0% 0% 0% 3 160 10 0% 0% 99% 1% 0% 0%

86% 14% 0% 0% 0% 0% 4 43% 0% 43% 0% 0% 14%

25% 75% 0% 0% 0% 0% 5 20% 40% 20% 20% 0% 0%

33% 33% 33% 0% 0% 0% 6 0% 4% 66% 27% 1% 2%

160 60 8% 8% 83% 0% 0% 0% 7 0% 1% 92% 7% 0% 0%

8 50% 50% 0% 0% 0% 0%

203

0% 6% 50% 41% 1% 2%

86% 14% 0% 0% 0% 0% 9 43% 0% 43% 0% 0% 14%

25% 75% 0% 0% 0% 0% 10 20% 40% 20% 20% 0% 0%

33% 33% 33% 0% 0% 0% 11 0% 4% 67% 28% 0% 2%

8% 8% 83% 0% 0% 0% 12 0% 1% 92% 7% 0% 0%

50% 50% 0% 0% 0% 0% 13 120 60 0% 5% 50% 43% 1% 2%

86% 14% 0% 0% 0% 0% 14 43% 0% 43% 0% 0% 14%

25% 75% 0% 0% 0% 0% 15 20% 40% 20% 20% 0% 0%

33% 33% 33% 0% 0% 0% 16 0% 3% 66% 28% 0% 2%

8% 8% 83% 0% 0% 0% 17 0% 1% 92% 7% 0% 0% 80 60 18 50% 50% 0% 0% 0% 0% 0% 4% 50% 43% 1% 2%

86% 14% 0% 0% 0% 0% 19 43% 0% 43% 0% 0% 14%

25% 75% 0% 0% 0% 0% 20 20% 40% 20% 20% 0% 0%

204

The following figures correspond to Entries 2 – 5 of Table 4.3 in Chapter 4. Table 4.3. Entry 2.

Figure S3. Reaction profile of the synthesised 85 wt% Ni/Al2O3/SiO2 catalyst. (38 –

90 µm).

Table 4.3. Entry 3.

Figure S4. Reaction profile of the synthesised and reduced 85 wt% Ni/Al2O3/SiO2 catalyst (38 – 90 µm).

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Table 4.3. Entry 4.

Figure S5. Reaction profile of the synthesised 85 wt% Ni/Al2O3/SiO2 catalyst (90 –

250 µm).

Table 4.3. Entry 5.

Figure S6. Reaction profile of the synthesised and reduced 85 wt% Ni/Al2O3/SiO2 catalyst (90 – 250 µm).

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S3: Runge-Kutta method (RK4) used in Berkeley Madonna Curve Fitting

The kinetic studies of batch reactions involved the curve fitting of experimental data to obtain kinetic constants of the N-alkylation of NH3 with BnOH using Berkeley

Madonna. The Runge-Kutta method (fourth-order formula, RK4) was chosen as the numerical solver because it generally gives less error than the Euler method and is a good compromise between computational cost and order of accuracy. The Runge-Kutta method is a family of methods that numerically integrates ordinary differential equations by using a trial step at the midpoint of an interval to cancel out lower error terms. [168] The fourth-order formula is as follows:

If y is a function of t, and

y(푡0) = 푦0 y is an unknown function of time t, which we wish to approximate. The rate at which y changes is a function of t AND y. At initial time (t0) the y value is y0. The function f and initial conditions, t0 and y0 are given.

If step size h>0, the approximation would be

ℎ 푦 = 푦 + (푘 + 2푘 + 2푘 +푘 ), 푛+1 푛 6 1 2 3 4

푡푛+1 = 푡푛 + ℎ

For n = 0, 1, 2, 3….. where the k values are defined as:

푘1 = 푓(푡푛, 푦푛)

ℎ 푘 푘 = 푓 (푡 + , 푦 + ℎ 1) 2 푛 2 푛 2

ℎ 푘 푘 = 푓 (푡 + , 푦 + ℎ 2) 3 푛 2 푛 2

207

푘4 = 푓(푡푛 + ℎ, 푦푛 + ℎ푘3)

yn+1 is determined by the present value (yn) plus the weighted average of four increments

(k1 to k4). These increments are the products of the sizes of the interval, h, and an estimated slope specified by function f.

k1 is the increment based on the slope at the beginning of the interval, using y (Euler's method)

푘1 k2 is the increment based on the slope at the midpoint of the interval, using 푦 + 2

푘2 k3 is again the increment based on the slope at the midpoint of the interval, using 푦 + 2

k4 is the increment based on the slope at the end of the interval, using 푦 + 푘3.

And greater weight is given to the midpoint increments when averaging.

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S4.1 Differential equations used in the Berkeley Madonna curve fitting simulation

The alkylation of ammonia with benzyl alcohol was conducted in batch. The model used for the simulation is as follows

Scheme S2. Three-reaction model of the N-alkylation of NH3 with BnOH with labelled kinetic constants.

The differential equations used in the model are shown below. It should be noted that

[NH3] and [H2O] were not available as experimental data, but were applied as initial conditions.

흏[푩풏푶푯] = −푘 [푁퐻 ][퐵푛푂퐻] + 푘 [퐵푛푁퐻 ][퐻 푂] 흏풕 1푓 3 1푟 2 2

− 푘2푓[퐵푛푁퐻2][퐵푛푂퐻]+푘2푟[퐵푛2푁퐻][퐻2푂]

− 푘3푓[퐵푛2푁퐻][퐵푛푂퐻]+푘3푟[퐵푛3푁][퐻2푂]

흏[푵푯 ] ퟑ = −푘 [푁퐻 ][퐵푛푂퐻] + 푘 [퐵푛푁퐻 ][퐻 푂] 흏풕 1푓 3 1푟 2 2

흏[푩풏푵푯 ] ퟐ = 푘 [푁퐻 ][퐵푛푂퐻] − 푘 [퐵푛푁퐻 ][퐻 푂] − 푘 [퐵푛푁퐻 ][퐵푛푂퐻] 흏풕 1푓 3 1푟 2 2 2푓 2

+ 푘2푟[퐵푛2푁퐻][퐻2푂]

흏[푩풏ퟐ푵푯] 흏풕

= 푘2푓[퐵푛푁퐻2][퐵푛푂퐻]−푘2푟[퐵푛2푁퐻][퐻2푂]−푘3푓[퐵푛2푁퐻][퐵푛푂퐻]+푘3푟[퐵푛3푁][퐻2푂]

흏[푵푯 ] ퟑ = 푘 [퐵푛 푁퐻][퐵푛푂퐻]−푘 [퐵푛 푁][퐻 푂] 흏풕 3푓 2 3푟 3 2

209

흏[푯 푶] ퟐ = 푘 [푁퐻 ][퐵푛푂퐻] − 푘 [퐵푛푁퐻 ][퐻 푂] 흏풕 1푓 3 1푟 2 2

+ 푘2푓[퐵푛푁퐻2][퐵푛푂퐻]−푘2푟[퐵푛2푁퐻][퐻2푂]

−휕[퐵푛푂퐻] + 푘 [퐵푛 푁퐻][퐵푛푂퐻]−푘 [퐵푛 푁][퐻 푂] = 3푓 2 3푟 3 2 휕푡

All three steps occur on the catalyst, but Ni was not included in this model as the number of catalyst sites is assumed to be constant, and is therefore part of the rate constants.

210

S4.2 Differential equations used in the Modified model in the Berkeley Madonna curve fitting simulation

The differential equations used in the modified model are shown below. It should be noted that [NH3] and [H2O] were not available as experimental data, but were applied as initial conditions.

Scheme S3. Four-reaction model of the N-alkylation of NH3 with BnOH with labelled kinetic constants.

흏[푩풏푶푯] 흏[푵푯 ] 흏[푵푯 ] 흏[푯 푶] , ퟑ , ퟑ & ퟐ remained the same with changes to the following 흏풕 흏풕 흏풕 흏풕 equations:

흏[푩풏푵푯 ] ퟐ = 푘 [푁퐻 ][퐵푛푂퐻] − 푘 [퐵푛푁퐻 ][퐻 푂] − 푘 [퐵푛푁퐻 ][퐵푛푂퐻] 흏풕 1푓 3 1푟 2 2 2푓 2

+ 푘2푟[푃ℎ퐶퐻푁퐵푛][퐻2푂]

흏[푷풉푪푯푵푩풏] = −푘′ [퐵푛푁퐻 ] + 푘′ [푃ℎ퐶퐻푁퐵푛] 흏풕 2푓 2 2푟

흏[푩풏ퟐ푵푯] 흏풕

′ = 푘 2푟[푃ℎ퐶퐻푁퐵푛]−푘2푟[퐵푛2푁퐻]−푘3푓[퐵푛2푁퐻][퐵푛푂퐻]+푘3푟[퐵푛3푁][퐻2푂]

All four steps occur on the catalyst, but Ni was not included in this model as the number of catalyst sites is assumed to be constant, and is therefore part of the rate constants.

211

S4.3 General Solver in kinetic model

A general solver is calculated by considering different data sets and fitting them with a universal set of kinetic values. The data sets may have different initial concentrations but must have the same reaction conditions.

Table S2. Sample code for calculating the general solver on Berkeley Madonna.

aCode Function

k1 = 1 Defines the rate constants

k1r = 0 (the initial values are

k2 = 1 chosen arbitrarily).

k2r = 0

k3 = 1

k3r =0

k4 = 1

k4r = 0

INIT A_XX = 0.596 Defines the initial

INIT B_XX = 1.2 concentrations of reagents

INIT C_XX = 0 A – G

INIT D_XX = 5.4

INIT E_XX = 0

INIT F_XX = 0

INIT G_XX = 0

RXN1_XX = k1 * A_XX * B_XX - k1r * C_XX * Defines each step of the

D_XX mechanism in terms of the

RXN2_XX = k2 * A_XX * C_XX - k2r * D_XX * reagent concentrations.

E_XX

RXN3_XX = k3 * E_XX - k3r * F_XX

212

RXN4_XX = k4 * A_XX * F_XX - k4r * G_XX *

D_1

d/dt(A_1) = -RXN1_XX - RXN2_ XX - RXN4_ XX Defines the change in

d/dt(B_1) = -RXN1_ XX reagent concentrations

d/dt(C_1) = RXN1_ XX - RXN2_ XX with time with differential

d/dt(D_1) = RXN1_ XX + RXN2_ XX + RXN4_ XX equations, which are

d/dt(E_1) = - RXN3_ XX + RXN2_ XX solved by the software.

d/dt(F_1) = RXN3_ XX - RXN4_ XX

d/dt(G_1) = RXN4_ XX aXX represents a unique identifier given to each dataset.

Crucially, the rate constants do not contain a unique identifier and so are common to every set of differential equations. Experimental data sets were used to determine the value of 6 variables (rate constants k1, k2, k3, k4; rate constants k2r and k3r were set to zero). The differential equation solver, Berkeley Madonna (www.Berkeleymadonna

.com), was used to calculate the model.

213

S5. Curve fits using reaction kinetics

Table 5.5 Entry 1. BnOH+NH3+H2O.

Figure S8. Four-reaction model fit for BnOH+NH3+H2O, Table 5.6 Entry 1.

214

Table 5.5, Entry 2. BnNH2+H2O.

Figure S9. Three-reaction model fit for BnNH2+H2O, Table 5.5 Entry 2.

Figure S10. Four-reaction model fit for BnNH2+H2O, Table 5.6 Entry 2.

215

Table 5.5, Entry 3. Bn2NH+H2O.

Figure S11. Three-reaction model fit for Bn2NH+H2O, Table 5.5 Entry 3.

Figure S12. Four-reaction model fit for Bn2NH+H2O, Table 5.6 Entry 3.

216

Table 5.5, Entry 4. Bn3N + H2O

Figure S13. Three-reaction model fit for Bn3N+H2O, Table 5.5 Entry 4.

Figure S14. Four-reaction model fit for Bn3N+H2O, Table 5.6. Entry 4.

217

Table 5.8. Entry 1 BnOH:NH3:H2O = 1:1:5

Figure S15. Two-reaction model fit for BnOH:NH3:H2O = 1:1:5

Figure S16. Three-reaction model fit for BnOH:NH3:H2O = 1:1:5, Table 5.8. Entry 1.

218

Table 5.8. Entry 2 BnOH:NH3:H2O = 1:1:10

Figure S17. Two-reaction model fit for BnOH:NH3:H2O = 1:1:10

Figure S18. Two-reaction model fit for BnOH:NH3:H2O = 1:1:10. Table 5.8. Entry 2.

219

Table 5.8. Entry 3 BnOH:NH3:H2O = 1:2:10

Figure S19. Two-reaction model fit for BnOH:NH3:H2O = 1:2:10

Figure S20. Three-reaction model fit for BnOH:NH3:H2O = 1:2:10. Table 5.8. Entry

220

Table 5.8. Entry 4. BnOH:NH3:H2O = 1:2:20

3 reaction model

Figure S21. Three-reaction model fit for BnOH:NH3:H2O = 1:2:20. Table 5.8. Entry

221

Table 5.8. Entry 6. BnOH:NH3:H2O = 6:1:0 2 reaction model

Figure S22. Two-reaction model fit for BnOH:NH3:H2O = 6:1:0

3 reaction model

Figure S23. Three-reaction model fit for BnOH:NH3 = 6:1:. Table 5.8. Entry 6.

222

Table 5.6. Entry 7 BnOH:NH3 = 1:3 2-reaction model

Figure S24. Two-reaction model fit for BnOH:NH3 = 1:3.

223

Table 5.11. Entry 1 BnOH:NH3 = 1:3 at 80 °C

Figure S25. Two-reaction model fit for BnOH:NH3 = 1:3 at 80 °C

Table 5.11. Entry 2 BnOH:NH3 = 1:3 at 120 °C

Figure S26. Two-reaction model fit for BnOH:NH3 = 1:3 at 120 °C

224

Table 5.11. Entry 3 BnOH:NH3 = 1:3 at 160 °C

Figure S27. Two-reaction model fit for BnOH:NH3 = 1:3 at 160 °C

Universal kinetic constants fit back into the starting Ni experiment.

Figure S28. Universal kinetic constants fit back into the starting Ni experiment.

225

S6 Characterisations

NMR

Benzyl amine 1H NMR (300 MHz, CDCl3): δ = 1.55 (s, 2H), 3.90 (s, 2H), 7.33 (m, 5H)

[169]

13C δ = 143.33, 128.55, 127.08, 77.40, 46.54 ppm[170]

GC-MS (EI)

Benzyl amine: m/z (%): 106 (100), 91(18), 79(39), 55(9)

2-Phenyl ethylamine: m/z (%): 121(50), 103(8), 91(100), 77(18), 65(52), 51(30)

2-Aminomethyl pyridine (2-picolyl amine) m/z: 106(82), 92(8), 79(100), 65(7),

52(30)

3-Aminomethyl pyridine (3-picolyl amine) m/z: 107(24), 92(7), 79(100), 65(7),

52(10)

Octan-2-amine: m/z: 128(2), 114(4), 84(2), 72(6), 56(10), 44(100)

2-methoxybenzyl amine: m/z: 136(100), 121(48), 106(52), 91(50), 77(50), 65(30),

51(25)

3-methoxybenzyl amine: m/z: 136(100), 121(18), 106(40), 91(28), 77(30), 65(18),

51(10)

226

TEM

The TEM images suggest that the particles are amorphous. The particle sizes were determined by 182 particles using four images. The particle size distribution is represented by a bar chart in Figure S29. The mean particle size was 4.66 nm. This was confirmed by XRD particle size estimations.

Particle size distribution 60 50 40 30

Counts 20 10 0 2 4 6 8 10 Size (nm)

Figure S29. TEM images of 65 wt% Ni-Al2O3/SiO2. The bar chart shows catalyst particle size distribution, data collected by measuring 182 particles using four images

227

Figure S30. TEM image of 85 wt%Ni/Al2O3/SiO2

228

XRD

Ni Ni (111) SiO SiO

Mt Mt Mt Mt

Spent Extrusion

Bentonite Intensity/a.u. Commercial

10 20 30 40 50 60 70

2/Degrees

Figure S31. XRD spectra for comparison of catalysts at different stages. Mt represents montmorillonite, a form of Al2O3/SiO2 present in bentonite clay

Particle sizes were estimated using the Scherrer equation, D=Kλ/βcosθ, as described in

Chapter 4. Using the commercial Ni (111) peak, a grain size of 4.2 nm was estimated.

This matches with the particle size distribution from TEM, which was an average of

4.66 nm.

229

Figure S32. XRD patterns of synthesised 85 wt% Ni/Al2O3/SiO3 catalysts. The graphs are identical. The bottom graph is labelled for clarity.

230

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