i .

THE COPPER-CATALYSED DEHYDROGENATION OF METHANOL

by

STEPHEN PATRICK TONNER B.Sc.(Hons.)

A Dissertation submitted to the School of Chemical Engineering and Industrial Chemistry in partial fulfilment of the requirements for the degree of Doctor of Philosophy.

The University of New South Wales May 1984 ii.

Doctor of Philosophy (1984) University of New South Wales (School of Chemical Engineering Sydney, Australia and Industrial Chemistry)

TITLE: The Copper-Catalysed Dehydrogenation of Methanol AUTHOR: Stephen Patrick TONNER, B.Sc. (Hons.), U.N.S.W. SUPERVISORS: Professor D.L. Trinm Associate Professor M.S. Wainwright NO. OF PAGES: i-xvi; 1-295. RESEARCH PUBLICATIONS: 1. Tonner, S.P., Wainwright, M.S., Trimm, D.L., and Cant, N.W., J. Chem. Eng. Data, 28, 59 (1983). 2. Tonner, S.P., Trinun, D.L., Wainwright, M.S., and Cant, N.W., J. Mol. Catal., 18, 215 (1983). 3. Evans, J.W., Tonner, S.P., Wainwright, M.S., Trimn, D.L., and Cant, N.W., Proceedings of the 11th Australian Chemical Engineering Conference, Brisbane, 509, Sept. (1983). 4. Tonner, S.P., Wainwright, M.S., Trimm, D.L., and Cant, N.W., Appl. Catal., In Press. 5. Tonner, S.P., Trinm, D.L., Wainwright, M.S., and Cant, N.W., Ind. Eng. Chem. Prod. Res. Dev., In Press. 6. Cant, N.W., Tonner, S.P., Trinm, D.L., and Wainwright, M.S., Submitted to J. Catal., April (1984). iii.

CANDIDATES CERTIFICATE

This is to certify that the work presented in this thesis was carried out in The School of Chemical Engineering and Industrial Chemistry University of New South Wales, and has not been submitted previously to any other university or technical institution for a degree or award.

Stephen Patrick TONNER B.Sc. {Hons.), U.N.S.W. iv.

ABSTRACT

Copper catalysts have been shown to be highly active and selective for the dehydrogenation of methanol to methyl formate at atmospheric pressure, and over the temperature range 180-240°C.

Copper chromite catalysts, which were particularly effec­ tive for the reaction, were characterized in detail in order to determine the mechanism of catalyst reduction, and to understand the nature of the active surface. Elemental analysis, surface area determination, X-ray diffraction and thermal gravimetric analysis were used to show that the active catalyst consisted of copper crystallites supported on chromia or cuprous chromite. Differences in activity, selectivity and stability amongst the copper chromite catalysts were related to the extent of catalyst reduction, and the degree of copper dispersion.

The specific activity of the copper chromite catalysts was low when compared with that for silica-supported and pure copper. Mass transfer effects and overestimation of metal surface area were ruled out, and the result was attributed to electronic interaction between copper and copper-chromium oxides. In the absence of support interaction, dehydrogenation activity was simply proportional to the available copper surface area. Differences in selectivity amongst various copper catalysts were related to the activity of the support phase for the decarbony­ lation of methyl formate.

The effectiveness of Raney copper catalysts was also examined, but despite exhibiting high initial activity, these v. catalysts underwent dramatic deactivation. To explain this phenomenon, a model was proposed involving the blockage of active surface within the catalyst pores by the polymerization of formaldehyde, an intermediate in the reaction mechanism.

Experiments with deuterium, isotopically labelled methanol and formaldhyde were carried out to investigate the reaction mechanism in detail. The very large kinetic isotope effect that was observed was consistent with the dehydrogenation of adsorbed methoxide to formaldehyde as being the rate controlling step. A thermodynamic isotope effect, that served to decrease the concentration of surface methoxide, was also identified. Identification of the mechanism whereby fonnaldehyde was converted to methyl formate, was not possible due to transesterification. vi.

ACKNOWLEDGEMENTS

This thesis could not have been completed without the assistance and co-operation of a great number of people. I would particularly like to express '1\Y gratitude to:

My supervisors, Professor David Trinn and Associate Professor Mark Wainwright, for their patience and comradeship, and for being consistently approachable.

Dr. Noel Cant for his valuable assistance with the isotope experiments, and for his contribution to all aspects of the research.

My fellow postgraduate students, especially John Evans and John Honig, for their good company. I wish them all the very best with their careers.

Ashley Deacon for his lively technical assistance.

Mrs. Vi Weatherill for sacrificing so nuch free time to type the thesis.

My parents, and all '1\Y family, for their understanding and encouragement over twenty five years.

To these and numerous other acquaintances, many thanks. vii.

TABLE OF CONTENTS

ABSTRACT iv ACKNOWLEDGEMENTS Vi TABLE OF CONTENTS vii

LIST OF TABLES xi LIST OF FIGURES xiii

CHAPTER 1 INTRODUCTION 1

2. LITERATURE REVIEW 4 2.1. Production of, and Use of Methyl Formate 4 2.1.1. Synthesis of methyl formate by carbonylation 4 2.1.2. Uses of methyl formate 9 2. 2. Reactions Invol v:ing the Dehydrogenation of Alcohols 14 2.2.1. Introduction to alcohol dehydrogenation 14 2.2.2. Dehydrogenation of ethanol 18 2.2.3. Dehydrogenation of methanol to formaldehyde 23 2.3. Aspects of Methyl Formate Synthesis by Dehydrogenation of Methanol 26 2.3.1. Introduction 26 2.3.2. Reaction mechanism 29 2.3.3. Steam reforming of methanol 36 2.3.4. Hydrogenolysis of methyl formate 39 2.4. Copper Chromite Catalysts 40 2.4.1. Introduction 40 2.4.2. The structure of copper chromite catalysts 44

3. PROJECT OBJECTIVES 52 viii.

TABLE OF CONTENTS (Cont'd.)

CHAPTER Page

4. EXPERIMENTAL TECHNIQUES 53 4.1. Catalyst Testing 53 4.1.1. Apparatus 53 4.1.2. Chromatographic analysis 57 4.1.3. Sources and purity of chemicals employed 59 4.1.4. Calculations 60 4.2. Catalysts 65 4.2.1. Copper chromite catalysts 65 4.2.2. Supported catalysts 67 4.2.3. Raney copper catalysts 68 4.2.4. Other catalysts 70 4.2.5. Catalyst pre-treatment 72 4.3. Catalyst Characterization 72 4.3.1. Atomic absorption analysis 72 4.3.2. Surface area characterization 4.3.2.1. Measurement of total area by nitrogen adsorption 73 4.3.2.2. Measurement of copper area by reaction with nitrous oxide 76 4.3.2.3. Single-point nitrogen adsorption 80 4.3.3. X-ray diffraction 81 4.3.4. Thermal gravimetric analysis 83

5 INVESTIGATION OF COPPER CHROMITE CATALYSTS 86 5.1. Introduction 86 5.2. Catalyst Characterization 88 5.2.1. Elemental Analysis 88 5.2.2. Measurement of total surface area and pore size distribution 90 5.2.3. Copper surface areas 91 5.2.4. X-ray diffraction 94 5.2.5. Thermal gravimetric analysis 102 5.3. Discussion 106 5.4. Conclusions 111 ix.

TABLE OF CONTENTS (Cont'd.)

CHAPTER Page

6 PRELIMINARY INVESTIGATIONS 112 6.1. Introduction 112 6.2. Thermodynamic Yields of Methyl Formate 112 6.3. Results 118 6.3.1. Activity and selectivity of copper chromite catalysts 118 6.3.2. Catalyst stability 122 6.4. Discussion 125 6.5. Conclusions 135

7 PROPERTIES OF COPPER-BASED CATALYSTS FOR METHANOL DEHYDROGENATION 136 7.1. Introduction 136 7.2. Preparation and Characterization of Catalysts 137 7.2.1. Copper powder 137 7.2.2. Copper supported on silica 137 7.2.2.1. . Impregnated catalysts 137 7.2.2.2. Ion exchanged catalysts 140 7.2.3. Copper supported on alumina 143 7.2.4. Copper supported on magnesia 145 7.2.5. Copper supported on chromia 146 7.3. Preliminary Investigations 147 7.4. Results 152 7.5. Discussion 154 7.6. Catalyst Selection 171 7.7. Conclusions 172 x.

TABLE OF CONTENTS (Cont'd.)

CHAPTER Page

8 BEHAVIOUR OF RANEY COPPER CATALYSTS 174

8.1. Introduction 174 8.2. Results 176 8.3. Discussion 188 8.4. Conclusions 200

9 REACTION MECHANISM 202

9.1. Introduction 202 9.2. Isotope Effects in Catalysis 208 9.3. Experimental 211 9.4. Isotope Experiments 216 , 9. 4.1. Results 216 9.4.2. Discussion 228 9.5. Experiments with Formaldehyde 234

9.5.1. Results 234 9.5.2. Discussion 237

9.6. Conclusions 241

10 CONCLUSIONS AND RECOMMENDATIONS 243

REFERENCES 248

APPENDICES 265 xi.

LIST OF TABLES

TABLE Page

2.1. Calculated equilibrium conversions of formaldehyde to methyl formate by 2HCHO ~ HCOOCH3 27 4.1. Retention times for gas chromatographic analysis 58 4.2. Relative molar responses (to methanol) 59 4.3. Sources and purity of chemicals employed 59 4.4. Suppliers and nominal compositions of copper chromite catalysts 66 4.5. Schedule for caustic soda addition in the preparation of Raney catalysts 69 4.6. Characteristics of Raney catalysts 70

4.7. Miscellaneous catalysts 71 5.1. Characteristics of copper chromite catalysts 89 5.2. Changes in surface area and crystallite size on reduction 92 5.3. Effect of N20 sample size on copper surface area 93 5.4. Tabulated values for X~ray diffraction maxima 95 5.5. Observed and predicted weight losses due to catalyst reduction 104 5.6. Weight loss by TGA for copper chromite catalysts 104 5.7. Predicted and observed weight losses for catalyst reduction at 220°C 106 5.8. Composition of copper chromite catalysts after reduction 110 6.1. Activation/deactivation of catalyst 0203 124 7.1. Characteristics of copper/silica catalysts prepared by impregnation 138 xii.

LIST OF TABLES (Cont 1 d.)

TABLE Page

7.2. Properties of reduced catalyst samples 148 7.3. Effect of methanol concentration on reaction rate as predicted from the kinetic expression of Myazaki and Yasumori [122] 164 7.4. Kinetic parameters for the dehydrogenation of methanol 166 7.5. Decarbonylation activity of support materials. 30 cm3min-l HCOOCH3 at 220°C 170 7.6. Rates of dehydrogenation on the more active catalysts 172 8.1. Characteristics of Raney catalysts 175

8.2. Evaluation of t/Rp 194 9.1. Results of CH 30D dehydrogenation 217 9.2. Results of CH30H/H2tD2 exchange 219 9.3. Results of CH30H/D2 exchange at 180°C 222 9.4. Kinetic isotope experiments at 180°C 224

9.5. Results of 11 CD30H/HCOOCH311 experiment 225 9.6. Results of 11 CD30D/CH3ow experiment 227 xiii.

LI ST OF FI GU RES

FIGURE

2.1. Flow diagram for methyl formate synthesis in formic acid production 8 2.2. Flow diagram of proposed two stage methanol synthesis 10 2.3. Uses of methyl formate 12 2.4. Mechanism of the Tischtschenko reaction 33 2.5. Mechanism of formaldehyde condensation 34 2.6. Hemi acetal mechanism for the synthesis of methyl fonnate 36 4.1. Experimental Apparatus for catalyst testing 54

4.2. Graphical determination of initial rate 64 4.3. Experimental apparatus for the measurement of copper surface area 78 5.1. Pore size distributions of copper chromite catalysts 92 5.2. XRD spectra for laboratory-prepared copper chromite (left) and Cucr2o4 (right) 96 5.3. XRD spectra for unreduced copper chromite catalysts. Left to right, catalysts 1,2,3 98 5.4. XRD spectra for leached copper chromite catalysts. Catalysts 1 (left) and 2 (right) 99 5.5. XRD spectra for reduced copper chromite catalysts. Left to right, catalysts 1, 2, 3 101 5.6. Weight loss profiles for reduction of copper chromite catalysts. From top: Catalysts 1, 2, 3, 41 Cucr2o4 103 6.1. Effect of temperature on thennodynamic yi,eld. P = 1 atm, YcH OH = 1 113 3 6.2. Effect of pressure on thennodynamic yield 220°C, YcH OH= 1 115 3 xiv.

LIST OF FIGURES (Cont'd.)

FIGURE

6.3. Effect of methanol concentration on thermodynamic yield. 220°C, P = 1 atm 116 6.4. Effect of hydrogen concentration and methyl formate concentration on thermodynamic yield. 220°c, P = 1 atm 117 6.5. Effect of LHSV on conversion and selectivity at 200°C 119 6.6. Effect of LHSV on conversion and selectivity at 220°C 120 6.7. Effect of LHSV on conversion and selectivity at 240°C 121 6.8. Stability of copper chromite catalysts. 220°C, LHSV = 26 h-1 123 6.9. Arrhenius plot for copper chromite catalysts 126 6.10. Effect of temperature on conversion LHSV = 25 h-l 127 6.11. Effect of temperature on selectivity. LHSV = 25 h-1 . 130 6.12. Decarbonylation activity of copper chromite 1808. YHCOOCH = 1, LHSV = 13 h-1 132 3

7 .1. Pore size distributions of copper/silica catalysts 139 7.2. Effect of LHSV on product distribution for copper alumina. 220°C 151 7.3. Deactivation of copper powder and copper chromia. 220°C. LHSV = 26 h-1 153 7.4. Effect of LHSV on conversion for copper chromite catalysts at 220°C 155 7.5. Effect of LHSV on selectivity for copper chromite catalysts at 220°C 156 7.6. Effect of LHSV on conversion for copper/ silica catalysts at 220°C 157 xv.

LIST OF FIGURES {Cont'd.)

FIGURE Page

7.7. Effect of LHSV on selectivity for copper/silica catalysts at 220°C 158 7.8. Effect of LHSV on conversion for copper/ silica {ion exchange) and copper/magnesia catalysts at 220°c 159 7.9. Effect of LHSV on selectivity for copper/ silica {ion exchange) and copper/magnesia catalysts at 220°c 160 7.10. Effect of LHSV on selectivity and conversion for copper/chromia catalyst at 220°c 161 7.11. The relationships between conversion and selectivity for copper catalysts at 220°C 162 7.12. Effect of methanol concentration on conversion for copper catalysts. 220°c, LHSV = 53 h-1 163 8.1. Deactivation profili of Raney copper. 220°C, LHSV = 26 h- 177 8.2. Concentration of formaldehyde during deactivation of Raney copper. 220°c, LHSV = 26 h-1 179 8.3. Reactivation profile of Raney copper at 220°C and 100 cm3min-l H2. Activity tested at 220°C, LHSV = 26 h-1 179 8.4. Effect of methanol concentration and space velocity on deactivation of Raney copper at 220°c 180 8.5. Effect of temperature on deactivation of Raney copper 181 8.6. Deactivation profiles of Raney copper catalysts prepared under different leaching conditions. 220°C, LHSV = 26 h-1 183 8.7. Deactivation profiles of Raney catalysts prepared froT various alloys. 220°c, LHSV = 26 h- 184 xvi.

LIST OF FIGURES (Cont'd.)

FIGURE

8.8. Effect of CO on deactivation of Raney 186 copper. 220°c; LHSV = 13 h-1; .A YCH OH = l; . 3 • YCH OH= 0. 5; 3 8 YCH OH= 0.5, Yeo= 0.2 3 8.9. Effect of water concentration on 187 deactivation of Raney copper. 220°C, LHSV = 26 h-1 • YH O < 0. 002; a YH O = 0. 05 ;& YH O = 0. 01 2 2 2 8.10. Mechanism for the deactivation of Raney copper catalysts 190 8.11. Visualization of the surface of a deactivated catalyst 197 8.12. The relationships between selectivity and conversion for Raney copper and copper/ silica catalysts. 220°C 199 9.1. Dimerisation mechanism for methyl formate formation 205 9.2. Hemi-acetal mechanism for methyl formate formation 207 9.3. 1H n.m.r. spectrum (Table 9.5) 214 9.4. 2 H n.m.r. spectrum (Table 9.5) 215

9.5. Product distribution for HCHO/He over 1 copper chromite 1808. 180°C; 300 cm3min- ; YHCHO = 0.3; 1 g catalyst 235 9.6. Product distribution for HCHO/He over Raney copper. 180°C; 300 cm3min-1; YHCHO = 0.3; 1 g catalyst 236 1.

CHAPTER 1

INTRODUCTION

Dramatic increases in the cost of oil over the last decade, along with the environmental and economic desirability of diverse and renewable energy supplies, has led to a resurgence of interest

in 11 C1 chemistry". This loosely defined term includes all proces­ ses that are based on reactions of the carbon monoxide molecule such as the Fischer-Tropsch process which converts carbon monoxide and hydrogen to a wide range of hydrocarbons including the gaso­ line fraction, the Oxo process by which oxygenated hydrocarbons are produced by the reaction of olefins and carbon monoxide, carbonylation processes such as the synthesis of acetic acid from methanol, and the synthesis of methanol from carbon monoxide and hydrogen [1].

The great advantage of c1 chemistry lies in the diversity of possible sources of synthesis gas (carbon monoxide and hydrogen). Synthesis gas can be prepared from natural gas by steam reforming, from coal by gasification, and from a wide range of hydrocarbons by partial oxidation [2].

Because the technology for converting synthesis gas to methanol is well established [2,3], processes utilizing methanol

as a feedstock are also included under the label 11 c1 chemistry 11 • Of these processes, the most widely discussed at present is the Mobil Process for converting methanol to gasoline [4] though the use of methanol as a fuel itself, or as a blend with gasoline has 2.

also created much interest [5]. Methanol of course, already has considerable importance as a raw material in the synthesis of formaldehyde, (the basis of a large synthetic polymer industry), methyl halides and amines, methyl methacrylate, ethylene glycol ethers and acetic acid [2].

A lesser known but still important use of methanol is in the synthesis of methyl formate, an important intermediate in the pro­ duction of formic acid, formamide, N,N dimethyl formamide and hydrogen cyanide. Japanese authors have recently expressed consi­

derable interest in this particular field of c1 chemistry [6].

Methyl formate has traditionally been produced by the base catalysed carbonylation of methanol

(1. 1)

as a first step in the synthesis of formic acid and formamide. The process, first patented by BASF in 1925 [7], is operated in the liquid phase, typically at about 80°C and 45 atm, and employs dissolved sodium metal as a catalyst [8]. Although plants have been operating successfully for about fifty years, the carbonyla­ tion process suffers from several distinct disadvantages. Firstly, because the process operates at moderately high pressures, plant costs are high [6]. Further, although hydrogen and nitrogen can be tolerated, the carbon monoxide feed must be totally free of water, oxygen or carbon dioxide, or catalyst deactivation will result [9]. The purification procedures needed also add to the plant costs. As it is impossible to obtain completely dry carbon 3.

monoxide and methanol, catalyst deactivation is always a problem, especially as the deactivated catalyst {sodium formate) is insol­ uble and can foul heat transfer surfaces and pumps [10].

An alternative process for producing methyl formate is the dehydrogenation process, first reported by Mannich and Geilmann in 1916 [11].

{1.2)

Despite the early discovery of the reaction, and the considerable advantages of a single reactant and low pressure operation, a pro­ cess has never been commercialized and research has largely been confined to compounding of catalysts in Patent literature [6].

It is the purpose of this thesis to compare the activity of a wide range of copper catalysts for the dehydrogenation in order to gain some insight into the nature of the catalytic reaction and to confidently propose the most effective catalyst for further study. A detailed investigation of the "copper chromite" type catalysts, first reported by Adkins et al. [12], is also undertaken in order to properly understand the nature of this particularly effective and interesting catalyst. Experiments have also been carried out with methanol isotopes and deuterium so as to better understand the mechanism of the dehydrogenation reaction.

As part of this Ph.D. research program an investigation of the kinetics of the base-catalysed carbonylation route to methyl formate was undertaken. The results of this study are embodied in two papers which are attached to the thesis as Appendix A. 4.

CHAPTER 2

LITERATURE REVIEW

2.1. Production of, and Use of Methyl Fonnate

2.1.1. Synthesis of methyl fonnate by carbonylation

Methyl formate is most commonly produced by the base­ catalysed carbonylation of methanol, a liquid phase reaction where methanol forms the reaction medium. The reaction, which is given by

{1.1) was first reported in 1856 by Berthelot [13] and later studied in more detail by Stahler [14], and Christiansen and Gjaldbaek [15]. The catalyst for the reaction is an alkali alcoholate {or alkoxide), fanned by dissolving an alkali metal, most commonly sodium, in methanol.

{2.1)

As the carbonylation is highly exothennic, low temperatures {70-100°C) are used to maximize the equilibrium yield. Early patents [7,16] used very high pressures (200-1000 atm) of carbon monoxide to ensure high rates of reaction. During the Second World War, a process was operated in Germany to produce methyl fonnate for conversion to hydrogen cyanide [17]. The synthesis was perfonned batchwise in autoclaves at 60-100°C and 10-200 atm Eighty percent methanol conversion was obtained over four hours. 5.

Hatano and Ando [18] have made a detailed study of this process investigating the effects of catalyst concentration, temperature and pressure on the rate of reaction and yield of methyl formate, as well as the lifetime of the catalyst. They showed that above 80°C, catalyst deactivation occurred by the formation of sodium formate and dimethyl ether.

NaOCH3 + HCOOCH3 + HCOONa + CH30cH3 (2.2)

Below 100°C, the major mechanism of catalyst deactivation arises from the presence of trace water.

(2.3)

NaOH + HCOOCH3 + HCOONa + CH30H (2.4)

In carbonylation experiments at constant pressure, Tonner et al. have shown how yield of methyl formate can be limited by catalyst deactivation at high levels of conversion, especially when temperatures in excess of 70°C are used [19]. In a separate study, Tonner et al. compared the rates of carbonylation of methanol with those of a number of higher molecular weight aliphatic alcohols [20,21]. Initial rates at low alcohol conversion(< 5 percent) were found not to be controlled by mass transfer considerations, involving transfer of carbon monoxide from the gaseous to the dissolved phase, but were strongly related to the electron donating effects of substituents around the reactive 0-H group.

At 50°C, rate constants were found to increase in the order:- 6.

methanol, benzyl alcohol « ethanol< n-propanol < n-butanol << i-butanol, i-propanol, t-butanol << sec-butanol.

The bulk of published data on the carbonylation reaction involves improvements in reactor design to maximize gas-liquid contact, facilitate the removal of reaction products and to main­ tain the solubility of sodium formate to prevent it from coating heat transfer surfaces and fouling pumps. Patents have dealt with recycle reactors employing continuous product removal [22,23], optimization of catalyst concentration [24], and improvements in· gas-liquid reactor design [25,26,27]. Leuteritz et al. have shown how methanol conversions of up to 99 percent can be obtained using a loop reactor [10].

Imyanitov et al. proposed a heterogeneous process where sodium carbonate is the catalyst [28]. However, catalyst activity is very low, and very high pressures (up to 300 atm) have to be used to achieve even small yields. Other authors [29,30] have used sodium-graphite catalysts utilizing the catalytic properties of alkali metal intercalation compounds. Catalyst preparation is complex however, and the systems are easily poisoned by trace water.

The major industrial application of the carbonylation of methanol has been in the synthesis of formic acid. Aguila and Horlenko [8] describe in detail the three stage process whereby methanol is converted to fonnic acid via the synthesis of methyl formate and formamide.

1. CO + CH30H ::t: HCOOCH3 (1.1)

2. HCOOCH3 + NH3 ~ HCONH2 + CH30H (2.5) 7.

3. {2.6)

In the first stage (Figure 2.1) methanol {containing 2 weight percent dissolved sodium) and carbon monoxide are fed to a simple backmix reactor operated at 80°C and 45 atm. Carbon monoxide is continuously withdrawn from the top of the reactor and cooled to condense out methanol and methyl formate, which are fed to a gas­ liquid separator. Carbon monoxide is recycled to the reactor. Liquid rich in methyl formate is continuously withdrawn from the bottom of the reactor and freed of dissolved carbon monoxide in the gas-liquid separator. The off-gas from the separator is vented through an absorption tower to remove methanol and methyl formate vapours. The liquid from the gas-liquid separator is passed to a distillation unit where methyl formate is removed. After the removal of insoluble deactivated catalyst, the unreacted methanol­ catalyst mixture is recycled to the reactor. The conversion of methyl fonnate to formic acid will be described in the next section.

Recently, much interest has been shown in the carbonylation reaction as the first stage in a two stage alternative synthesis of methanol, first proposed by Christiansen (31].

{1.1)

{2.7)

The second stage involves hydrogenolysis of methyl fonnate and has been part of a detailed study by Evans et al. using copper catalysts [19,32,33,34].

co co

E'NDS E'NDS

TOWER TOWER

Catal',st Catal',st

blOwOOwn blOwOOwn

HEAVY HEAVY

Mttihyt Mttihyt

fu-rna~ fu-rna~

FORMATE FORMATE

TOWER TOWER

synthesis synthesis

METHYL METHYL

Mt,thanol Mt,thanol

formate formate

~ ~

t-aeyc:I~ t-aeyc:I~

methyl methyl

TANK TANK

production production

RUNDOWN RUNDOWN

CO CO

for for

acid acid

diagram diagram

formic formic

in in

Flow Flow

CAS-LIGUID CAS-LIGUID

SEPARATOR SEPARATOR

2.1. 2.1.

PUl"

LJ LJ

CQ CQ

FIGURE FIGURE

FORMATE FORMATE

REACTOR REACTOR

METHYL METHYL

co co

ME.~NOL ME.~NOL CATALYST CATALYST 9.

The overall reaction is the same as that for the conventional methanol synthesis carried out over Cu/Zn0/Al 2o3 catalysts [2,35].

(2.8)

Whereas conventional methanol synthesis is carried out at 250-300°C and 5,000-10,000 kPa, reaction 1.1 can be operated at 4000 kPa and 70°C [19], while reaction 2.7 proceeds with high yields and selecti­ vities at l60°C and 101 kPa over copper chromite and Raney copper catalysts [19,32,36]. The potential energy savings using the two stage process therefore are considerable.

Tonner and Evans [19] investigated the feasibility of the two stage process and their flow diagram for the proposed process is presented in Figure 2.2. The carbonylation is carried out using a simplified version of the process of Figure 2.1, and the hydrogen­ olysis of the methyl formate in a tubular reactor. Deactivation of the carbonylation and hydrogenolysis catalysts present problems in the implementation of this process along with the increase in engineering costs over the single stage process.

Brendlein has proposed a process whereby both reactions, 1.1 and 2.7, are carried out in the one vessel operated at 100-130°C and 30-60 atm [37]. A dispersion of copper-chromite is used to effect the hydrogenolysis step. Yields were not quantified however, and catalyst lifetime was not discussed. Recently a similar process, operating on a continuous basis, has been described [38].

2.1.2. Uses of methyl formate

Methyl formate is not an important industrial chemical in Backmixreactor CO Stripper MeOH & Hydrogenolysis Hydrogen Methanol recovery catalyst reactor stripper

recovery 1 H~O-purge

I • CO recycle I T T + MeF recycle

MeOH + CAT cb (0

MeOH M e()-1 product + and CAT. recycle recycle

..... FIGURE2.2. Flow diagram of proposed two stage methanol 0 synthesis 11.

its own right. It does however have considerable significance as an intermediate in the synthesis of a number of important chemical products. The industrially important reactions of methyl formate are represented in Figure 2.3.

The most important industrial use of methyl fonnate is in the synthesis of formic acid via formamide (Paths 1 and 3 in Figure 2.3). In Path 1, formamide is produced by reaction with liquid a111110nia at 2-6 atm and 80-100°C [39].

(2.5)

Formic acid is then produced by formamide hydrolysis with 68-74 percent sulphuric acid (Path 3) [40].

Formic acid, which is an important chemical in dyeing and painting, pharmaceuticals, tanning and silage preservation, is easily libera­ ted from the a11111onium sulphate, which can of course be used as a fertilizer [8]. About one third of the world formic acid production (about 400,000 tonnes p.a.) is carried out via the methyl formate­ formamide route [8,40]. Formic acid can also be produced by direct hydrolysis of methyl formate (Path 4) [41].

(2.9)

Removal of the formic acid from.the excess water is difficult but can be achieved by extractive distillation using N-formylmorpholine ..... N

7

C~COOH

16

CO

2

}

3

3

formate

Is

methyl

of

HCOOCH

Uses

HCON(CH

2.3.

4

FIGURE

3

HCOOH

2

1

2 '

HCN

HCONH 13. as a solvent, or by weak salt fonnation between formic acid and 1-pentylimidazole [8].

Formamide, produced via Path 1 in Figure 2.3 has use as a solvent and chemical intermediate in the synthesis of imidazoles, pyrimidine and triazines [39]. A potential use is in the synthesis of hydrogen cyanide by thermal decomposition of formamide at 500- 6000C over alumina according to the reaction [17,18,42]

(2.10)

Dimethyl formamide, an important polar solvent in the production of PVC and polyacrylonitrile, can be produced by the reaction of methyl formate and dimethylamine at 80-100°C (Path 5) [43]

(2.11)

High purity carbon monoxide can be produced by the decarbony­ lation of methyl formate which proceeds with high selectivity over basic oxides between 250 and 400°C, (Path 6) [6,44]

(2.12)

The synthesis of acetic acid by isomerisation of methyl formate has been proposed but no catalysts have been suggested (Path 7) [6,45]

(2.13) 14.

2.2. Reactions Involving the Dehydrogenation of Alcohols

2.2.1. Introduction to alcohol dehydrogenation

The catalytic dehydrogenation of alcohols is important in the synthesis of a wide range of aldehydes and ketones. Acetone, for instance, has long been produced by the dehydrogenation of iso­ propyl alcohol over zinc oxide or copper catalysts [46] according to

300-400°C -+ (2.14)

The reaction mechanism for alcohol dehydrogenation (to be discussed further in section 2.3.2) is believed to involve aldehyde formation by simultaneous removal of a methyl and hydroxyl hydrogen atom from the adsorbed alcohol [47,48].

(2.15)

The dehydrogenation of alcohols over copper metal has been shown to have a particularly low activation energy [47,49], and it is this catalyst that is generally used. The vast amount of patent literature that exists on the copper-catalysed dehydrogenation of c3+ alcohols (alcohols with more than three carbon atoms) for specialized organic synthesis cannot be covered here. However, a comprehensive review of the range of catalysts that have been used has been presented elsewhere [50].

Earliest studies of .the dehydrogenation of alcohols over copper were carried out by Palmer and Constable. Palmer [51,52], in studies of ethanol dehydrogenation over electrolytic copper and 15.

reduced copper oxide, first proposed that zero valent copper was the active site for dehydrogenation. In a later study [43], Palmer investigated the effects on activity of combining copper with various oxides. These results are clouded due to the lack of suit­ able catalyst characterisation techniques at the time of the study. Palmer and Constable observed that primary alcohols gave similar rates of dehydrogenation over reduced copper and deduced that the adsorption of the -CH20H group onto the catalyst surface was the initial step in the reaction [54]. Constable further proposed the removal of the hydroxyl hydrogen to be the rate controlling step and deduced that the copper surface was covered with a unimolecular film of adsorbed alcohol [55,56].

The preparation in 1931 by Adkins et al. [57] of a copper chromite catalyst provided catalyst researchers with a highly active copper catalyst with reproducible catalytic activity. While the main application of this catalyst has been in the hydrogenation of esters in specialized organic synthesis [58], considerable research has also been carried out into its use in dehydrogenation. Dunbar [59,60] investigated the dehydrogenation of a range of c3+ alcohols over copper chromite, and together with Adkins et al. [61] noted the superiority in activity and lifetime of the copper chro­ mite catalysts over other types of copper catalyst. Dunbar also made the first attempts to explain the formation of high molecular weight condensation products observed in the dehydrogenation of c3+ alcohols. He identified four important reactions which in the case of n-butanol dehydrogenation are 16.

{l) Dehydrogenation

{2.16)

(2) Dehydration

{2.17)

(3) Aldol condensation {followed by dehydration)

2n-C3H7CHO 4 CH3CH2CH2CH{OH)CH(C2H5)CHO {2.18)

4 CH3CH2CH2CH = C{C2H5)CHO {2.19)

(4) Tischtchenko reaction

{2.20)

In a later study, Ipatieff and Haensel [63] further emphasized the necessity for both dehydrogenation and dehydration activity to promote the Aldol condensation {Equation 2.20).

Isopropanol dehydrogenation in the temperature range 200- 3000C over copper catalysts prepared by alkali precipitation of

Cu{OH) 2, followed by reduction at 180°C, has been studied in much detail by Kawamoto [63,64,65]. Both dehydrogenation and condensa- tion activity were influenced by the nature of the alkali precipitant. When sodium and potassium bicarbonates were used, the evolution of carbon dioxide yielded a porous catalyst and dehydro­ genation activity was highest. When more alkaline precipitants such as 17. sodium and potassium hydroxides were used, condensation activity was maximized leading to the conclusion that condensation activity was related to the existence of unreduced copper, present due to the inhibiting effect of trace alkali on reduction. Catalytic dehydration activity was unimportant and the primary condensation product was thought to be fonned via an intermolecular dehydration reaction between isopropanol and the dehydrogenation product, acetone according to

(2.21)

These conclusions were also extended to the dehydrogenation of secondary butanol [66].

In the dehydrogenation of primary alcohols, Kawamoto also found that condensation was promoted by alkalinity [67]. The major condensation products, esters, were thought to be formed via a hemi acetal intermediate according to OH I RCH 0H + RCHO 4 R-C-OCH R (2.22) 2 I 2 H (hemi acetal) OH I

R-C-OCH R 4 R-C-OCH R (2.23) a 2 ., 2 H 0 (ester)

Brihta and Luetic [68] studied the dehydrogenation of c1-c4 alcohols over Raney copper at 150-250°C. Activation energies were 18.

found to increase in the order n-butan-2-ol < propan-2-ol << propan-1-ol < n-butan-1-ol < ethanol << methanol. These trends were readily explained by electron directing effects. Sznajder [69] used Raney copper to study the dehydrogenation of cyclohexanol. Deactivation was attributed to surface coverage with high molecular weight condensation products.

Whilst copper is the most active and selective catalyst for alcohol dehydrogenation, oxides possessing dehydrogenation activity have also been used. Komarewsky et al. have investigated the dehydrogenation and condensation activity of chromia for c4-c8 alcohols at 400-500°C [70,71]. Gershbein et al. have studied the dehydrogenation of isopropanol by magnesia at 490°C [72], and Adkins and Lazier the dehydrogenation of c2-c4 alcohols over zinc oxide at 350-450°C [73]. The temperatures required for catalyst activa­ tion are at least 100°C in excess of those required for copper catalysts and considerable dehydration is observed in all cases.

2.2.2. Dehydrogenation of ethanol

The dehydrogenation of ethanol over copper catalysts was used prior to the Second World War to produce acetaldehyde [74]

(2.24)

However, since that time the process has been superseded by the liquid phase oxidation of ethylene [74]. The reaction is discussed in some detail in this section because, of all the alcohol dehydro­ genations, it most closely resembles the dehydrogenation of methanol. The main similarities are:- 19.

1} Temperatures employed in ethanol dehydrogenation (250-300°C} are only slightly higher than those employed for the dehydrogenation of methanol to methyl formate (200-250°C}.

2} Copper-based catalysts are preferred for both reactions.

3} The primary ester (ethyl acetate or methyl formate} can be formed with high selectivity in both cases.

4} In ethanol and methanol dehydrogenation, deactiva­ tion is observed over the copper catalysts employed.

The pioneering studies by Palmer and Constable of copper catalysts using ethanol dehydrogenation as a model reaction, have been reviewed in the previous section (2.2.1}. Prior to these studies, Armstrong and Hilditch [75] had studied the hydrogenation of acetaldehyde and the dehydrogenation of ethanol over copper and nickel. Copper was found to be much more selective than nickel, which caused decomposition of acetaldehyde to carbon monoxide and methane, and water was found to improve the yield of acetaldehyde in the dehydrogenation. To explain the latter result, the displace­ ment by water of adsorbed acetaldehyde from the catalyst surface was suggested. Adkins and Lazier [76] coll11lented on the differences in activity of copper catalysts prepared by different methods and suggested that traces of unreduced copper oxides were responsible for dehydrogenation activity. However, differences in active metal surface area provide a more likely explanation. 20.

Dolgov et al. [77] first showed how ethyl acetate could be made the major product of ethanol dehydrogenation over copper-ceria catalysts at 275°C. In a subsequent study, Dolgov et al. [78] proposed the overall reaction to be:-

(2.24)

2CH3CHO ~ CH3COOCH2CH3 (2.25) (Ethyl Acetate)

The condensation reaction (2.25) was found to be favoured by high precipitant alkalinity in the catalyst preparation. This is in agreement with the.later results of Kawamoto [63-67] discussed in the previous section (2.2.1.). The addition of small amounts

(0.1-2 percent) of dehydrating agents (A1 2o3, Th02, Ce02 , Ti02 and

Cr2o3) was found to enhance dehydrogenation activity, though how much of this effect was due to increased surface areas is uncertain.

The formation of acetic acid, butanol and other c2-c6 condensation products was explained via acetaldehyde dehydration and recombina­ tion. Balandin et al. [79] proposed deactivation of ethanol dehydrogenation catalysts to be due to carbon formation derived from these condensation products.

Dolgov and Nizovkina [80] further investigated the dehydro­ genation of ethanol and proposed the initial dehydrogenation to acetaldehyde (Equation 2.24) to be rate determining. By passing acetaldehyde/ethanol mixtures over the copper catalyst, it was deduced that ethyl acetate was formed by the combination of two acetaldehyde molecules 21.

(2.25) rather than by reaction between ethanol and acetaldehyde.

(2.26)

Catalyst deactivation was attributed to polymerisation of the acetal­ dehyde intermediate.

Church and Joshi [81] used basic promoters (MgO, ZnO) to enhance ethyl acetate fonnation in the dehydrogenation of ethanol over supported and unsupported copper. The addition of chromia

(Cr2o3) was found to increase catalyst life. Shorter contact times than in previous studies were used in order to minimize condensation thus obtaining acetaldehyde as the major product.

The first comprehensive kinetic study of alcohol dehydrogena­ tion was carried out by Franckaerts and Froment [82] who studied ethanol dehydrogenation to acetaldehyde over the copper catalyst of Church and Joshi [81]. The initial rate approach was used to verify the rate expression

r = (2.27) where the subscripts A, R, S, Wrefer to ethanol, acetaldehyde, hydrogen and water respectively, and

k = rate coefficient (mole h- 1g.cat-1) -1 KN = adsorption coefficient of species N (atm ) 22.

PN = partial pressure of species N (atm)

K = equilibrium constant (atm)

The rate controlling step was deduced to be surface reaction on dual sites. Deactivation was attributed to acetaldehyde polymer­ isation. This approach was also used by Peloso et al. [87] and Venugopal et al. [83] using copper chromite and copper-asbestos catalysts respectively.

Prasad and Menon [84] also attributed deactivation of copper­ alumina to acetaldehyde polymerisation, as well as to the formation of high molecular weight condensation products favoured by acidic sites. Hydrogen, however, was believed to scavenge the surface of polymerisable products and inhibit deactivation. Alumina was shown to promote the condensation to ethyl acetate and copper-alumina co-sites were believed to be important in achieving high yields of ethyl acetate.

Sundaram and Ibrahim [85] showed how the addition of 0.8 per­ cent nickel to copper-pumice enhanced dehydrogenation activity by preferentially adsorbing hydrogen. Moorjani et al. [86] used chromia as a promoter to reduce deactivation and attributed the effect to an increase in the activity of the surface for the oxidation of carbonaceous deposits. Water was also used to scavenge the surface of polymerisable products.

Oxides with dehydrogenation activity have also been used to study ethanol dehydrogenation. Schwab et al. [88] investigated the dehydrogenation and dehydration activity of ZnO, Ti02, Cr2o3, Al 2o3, CaF2, CeO, Th02, C and Si02• Pandao et al. [89,90] have studied 23. the dehydrogenation of ethanol on chromia and found that the high temperatures required for catalyst activation(> 380°C) led to dehydration and acetone formation. The condensation reaction to ethyl acetate (Equation 2.25) was shown to take place on acidic dehydration sites indicating strong adsorption of acetaldehyde on these sites. Kuriacose and Sastri [91] have made a detailed study of the electronic nature of these sites. Takezawa et al. [92] have studied surface species involved in the dehydrogenation of ethanol on magnesium oxide. Again, high temperatures (> 350°C} were required for catalyst activation.

2.2.3. Dehydrogenation of methanol to formaldehyde

By analogy with the dehydrogenation of ethanol discussed in the previous section (2.2.2), the first step in the dehydrogenation of methanol to methyl formate should be the dehydrogenation of methanol to formaldehyde which takes place according to

(2.28)

For this reason, the reaction will be considered in some detail here.

The dehydrogenation of methanol has received much attention because of the potential advatanges it offers over the traditional synthesis of formaldehyde by methanol oxidation over silver oxide or iron-molybdenum oxide catalysts [93] by

(2.29) 24.

Considerable quantities of by-products such as carbon monox­ ide, carbon dioxide, methyl formate and formic acid, are formed in the oxidation process [93], the corrosive properties of formic acid being a particular problem [6]. As no oxygen is required in the dehydrogenation process, selectivity could be enhanced. Further, as no water is formed in Equation 2.28, formaldehyde purification is simplified. However, the dehydrogenation process is highly endothermic and so high temperatures (600-650°C) are required with the decomposition of formaldehyde

HCHO ~ CO+ H2 (2.30) likely to be significant [94,95]. The advantages and disadvantages of the oxidation and dehydrogenation processes have been discussed by Ghosh and Ghosh [95].

Considerable research has been carried out into preparing copper catalysts that are active, stable at high temperature and capable of supressing formaldehyde decomposition. Typically, zinc alloys or mixtures of copper and zinc oxides with selenium and tellurium oxides or sulphur as minor components, have been used. Chono and Yamamoto [6] have reviewed the available patent literature.

Copper catalysts have also been used by Punderson [96] to effect the oxidation of methanol to formaldehyde (Equation 2.29). Wachs and Madix [97] have investigated the surface species involved in the oxidation of methanol on copper 110.

If methanol is passed over copper at long contact times, complete decomposition to carbon monoxide and hydrogen by the reaction 25.

(2.31)

is observed. Ghosh and Chakravarty were first to attribute this decomposition to dehydrogenation of methanol followed by copper catalyzed decomposition of formaldehyde [98] according to

(2.28)

HCHO + CO+ H2 (2.30)

(2.31)

Catalyst deactivation was also reported, Ghosh and Baksi [99] improved the life and activity of copper powder for methanol decom­ position at 180-200°C by incorporation of 0.2 percent ceria or thoria in the catalyst. Poisoning of copper by chlorine and sulphur compounds was also reported [100].

Lawson and Thomson [101] used a circulating constant-volume apparatus to study the activity of copper oxide and copper wire for methanol decomposition at 300°C. Copper oxide, which was reduced during the course of reaction, was much mo~e active than copper wire, and was proposed to be the true active catalyst. The results of Ghosh et al. [98,99,100] who reported the activity of reduced copper, were attributed to traces of unreduced oxide. Surface area considerations were not investigated however.

Borisov et al. [102] showed the dehydrogenation to be rate determining in methanol decomposition on copper films at 290-400°C. Eversole [103] has patented a copper-nickel catalyst for methanol 26. decomposition at 350-400°C. A comprehensive study of the mechanism of methanol decomposition over nickel wire has been carried out by Yasumori et al. (117].

In the decomposition of methanol on zinc and magnesium oxides, formaldehyde is also considered to be an intermediate [104,105,106].

Aika et al. (107] have compared the selectivities of group VIII metals on various supports for the hydrogenation of formalde­ hyde by the reaction

{2.32) at 250°C. Platinum supported on zinc oxide gave the highest selec­ tivity {81 percent) while a Cu-Zn0-Al 2o3 methanol synthesis catalyst also proved highly effective.

2.3. Aspects of Methyl Formate Synthesis by Dehydrogenation of Methanol

2.3.1. Introduction

In the previous section {2.2.3) it was shown that formalde­ hyde could be formed as the major product of methanol dehydrogenation over copper catalysts at 600-650°C. In this respect, methanol and ethanol dehydrogenation were similar as acetaldehyde can be readily formed in ethanol dehydrogenation {Equation 2.24). Since ethyl acetate, formed from acetaldehyde condensation by Equation 2.25, can be produced as the major product of ethanol dehydrogenation, it would be expected that methyl fonnate could be obtained by the dehydrogenation of methanol followed by the condensation of formaldehyde according to 27.

(2.28)

2HCHO + HCOOCH3 (2.33)

(1.2)

This result is achieved by carrying out the dehydrogenation at lower temperatures (200-250°C) where the highly exothermic condensation reaction (Equation 2.33; 6H~98 = -118 kJ mol) proceeds to completion. In Table 2.1, equilibrium conversions and mole fractions of formaldehyde based on Equation 2.33 and thermodynamic data [108], are calculated for the temperature range 200-800°C. Clearly, formation of methyl formate at the expense of formaldehyde is overwhelmingly favoured below 400°C.

TABLE 2.1. Calculated equilibrium conversions of formaldehyde to methyl formate by 2HCHO + HCOOCH 3 Temperature Equilibrium Constant Conversion (oc) (atm) (%)

200 748886 100 300 3982 99.2 400 100.4 95.1 500 6.560 80.9 600 0.801 51.2 700 0.151 21.1 800 0.0387 7.0

(See Appendix II for calculations) 28.

Mannich and Geilmann [11] first observed the formation of methyl fonnate in the dehydrogenation of methanol over copper at 240-260°C. Ivannikov and Zherko [109] obtained methyl formate at 170-220°C over the copper catalysts used by Ghosh and Baksi to effect methanol decomposition at lower space velocities [99]. The dehydrogenation process was first patented by Willkie [110]. Charles and Robinet [111] patented a Raney copper catalyst for the reaction in 1950.

In the last decade, patent literature on the synthesis of methyl fonnate by dehydrogenation has concentrated on the prepara­ tion of copper catalysts by the decomposition of copper acetates, hydroxides, sulphates, nitrates and carbonates, to copper oxide followed by reduction to copper metal [112,113]. The addition of basic oxides such as CaO and MgO,Group IVb metal oxides such as

Zr02 and Ti02, oxides of the Lanthanides and Actinides [114], zinc oxide along with support materials such as silica, chromia and alumina, have all been claimed to improve activity and selectivity of the catalyst. Chono and Yamamoto [6] review the available patent literature. Although comparisons between patent catalysts are difficult due to differences in temperatures and space veloci­ ties employed, high selectivities (>85 percent) are generally evident.

de Pinillos and Victor [44] have patented a process for producing methyl formate with high selectivity using copper chromite and copper-zinc oxide catalysts. Morikawa et al. have shown the copper ion-exchanged fonn of FluQr.Tetra-Silisic Mica (Cu-TSM) to be highly active for methanol dehydrogenation to methyl fonnate [116]. 29.

High copper dispersion, and low support acidity of Cu-TSM was thought to be responsible for its superior activity and selectivity when compared with copper on silica catalysts and copper-exchanged zeolite catalysts [116].

2.3.2. Reaction mechanism

The mechanism of alcohol dehydrogenation on metals is not properly understood. Brihta and Leutic [68] first expressed the mechanism in terms of separate cleavage of the hydroxyl and a-C-H bonds which can be represented by

(2.34)

RCHO + H(a) (2.35)

(2.36)

It was later stated that steps 2.34 and 2.35 occurred simultaneously [47], a view supported by Miyamoto and Ogino in explaining ethanol dehydrogenation over molten indium [118]. Other authors have proposed removal of a-C-H (Equation 2.35) to be rate controlling in alcohol dehydrogenations over copper [119].

Yasumori et al. [120,121,122] have studied the decomposition and dehydrogenation of methanol on copper and nickel wires. The main products of methanol decomposition on copper at 280°C and 21.8 nm Hg were methyl formate and carbon monoxide. Formaldehyde was observed in the initial stage of reaction only. The rate equation was found to be 30.

19 - r0 = 5.5xl0 g x exp( 10,000/RT)P (2. 37) l+lxl0-3 exp (4600/RT)P which suggested the dehydrogenation of adsorbed methanol to formal­ dehyde to be rate determining as adsorption and kinetic terms for methanol only are apparent ( 11 r II is in molecules cm- 2sec-1; 11 g11 is . 0 a "roughness factor" of the surface; 11 P11 is partial pressure of methanol in mm Hg). Copper wire was found to be inactive for the exchange reaction of methanol with deuterium, suggesting that methanol was non-dissociatively adsorbed. Nickel wire, on the other hand, was active for the exchange, suggesting dissociative adsorp­ tion as CH30(a) (adsorbed methoxide). Methanol decomposition over nickel wire yielded only carbon monoxide and hydrogen, this diffe­ rence in selectivity when compared to copper, being related to decomposition of the adsorbed methoxide species.

(2.38)

(2.39)

(2.40)

(2.41)

Differences in the electronic configuration of copper and nickel were used to explain the differences in adsorbed states of methanol [122]. 31.

Borisov et al. [102] studied the adsorption isotherms of methanol and formaldehyde on copper film and confirmed the form of the rate equation (2.37) of Miyazaki and Yasumori [122]. Non­ dissociative adsorption of methanol was also proposed but the rate limiting step:-

(2.42) was suggested to occur as a series of steps with the elimination of one, then another hydrogen atom from the adsorbed methanol.

The full mechanism of Yasumori et al. [121,122] is shown below:-

(2.43)

(2.42)

(2.44)

(2.45)

HCOOCH3(a) + 2CO + 2H2 (2.46)

The formation of methyl formate by formaldehyde condensation in Equation 2.45 is consistent with the mechanism of ethyl acetate formation in ethanol dehydrogenation outlined in section 2.2.2. Ester formation from aldehydes over solid catalysts was first repor­ ted by Adkins et al. [123] who studied the condensation of 32.

acetaldehyde to ethyl acetate over zinc chromite at 360°C and 200 atmospheres. The formation of ethyl acetate was attributed to the Tischtschenko reaction first reported in 1906 [124].

In the Tischtschenko reaction, aldehydes are converted to esters by treatment with aluminium ethoxide

Al(0Et)3 > (2.47)

acetaldehyde ethyl acetate

The mechanism of the reaction is uncertain, but is likely to resem­ ble that in Figure 2.4 [125]. A close resemblance to the base­ catalysed Cannizaro reaction

(2.48) is noted [126]. Adkins et al. [123] proposed that zinc chromite, like aluminium ethoxide in Figure 2.4, was capable of activating aldehyde molecules for the Tischtschenko reaction. Adkins has also studied the reaction over a range of alkoxide catalysts.

Miyazaki and Yasumori [121,122] showed methyl formate to be the only product when fonnaldehyde was passed over copper wire at 150-180°C. Further, deuterium exchange experiments showed formal­ dehyde to be non-dissociatively adsorbed. The dimerization mechanism of Figure 2.5 was used to account for formaldehyde condensation to methyl fonnate on copper (Equation 2.45). Kotowski (127) has attributed methyl formate formation in methanol synthesis 33.

m m -4-w - -0 m , _ 4- t <(-- - -

FIGURE 2.4. Mecha~ism of the Tischtschenko reaction ?/l ,) ..

0 H--1H ': II 2 H- C- H + 2tt ~---__._~ C = 0 H·- C-H II 0 .• •.

H » H I i I H- · H-C-0 c==o• • • • I I H·-C-H H C-H II II 0 0 • ••

FIGURE 2.5. Mechanism of formaldehyde condensation 35. over Cu-Zn0-Al 203 catalysts to formaldehyde condensation and further, claimed the reaction to be favoured by catalyst alkalinity (128].

Ai, in attempting to explain the formation of methyl formate in methanol oxidation on Sn02-Mo03 catalysts, was first to propose an alternative mechanism involving reaction between methanol and formaldehyde according to

(2.49)

However, methanol was found not to increase the yield of methyl formate from formaldehyde, and the dimerization mechanism was accep­ ted (129]. Ai has further stressed the need for both acidic and basic catalyst functions to effect the dimerization (130].

Recently, Takahashi et al. (131] have studied the mechanism of methyl formate formation from formaldehyde over copper-silica catalysts at 100-180°C. At formaldehyde partial pressures of 0.01 atmospheres, the addition of methanol was found to dramatically increase the conversion to methyl formate. However, the possibility of methyl formate formation by methanol dehydrogenation (Equation 1.2) could not be discounted. When formaldehyde and deuterated methanol (CD30H) were passed over the catalysts, HCOOCD3 was pre­ dominantly formed. This result could not be explained by the dimerization mechanism of Figure 2.5 where formaldehyde loses a hydrogen atom to the methyl group of the resultant methyl formate. Instead, reaction between methanol and formaldehyde by a hemi acetal mechanism similar to that in Equations 2.22 and 2.23 was proposed (Figure 2.6). 36.

or

rhc..,OH H/ 'OCH 3 + HCOOCH 3

FIGURE 2.6. Hemi acetal mechanism for the synthesis of methyl formate

2.3.3. Steam reforming of methanol

The steam reforming of methanol to carbon dioxide and hydro- gen by

(2.50)

will be considered here as it is believed to proceed via methyl formate produced by methanol dehydrogenation. Catalysts active for the steam reforming reaction then, would also be likely to be active for the dehydrogenation (Equation 1.2). Steam reforming of methanol, which has been proposed as a means of producing pure hydrogen (133, 134,143]. was first patented by Prigent and Sugor [132). Pour et

al. (133] studied the reaction at 200-250°C over Cu0-Zn0-Cr2o3-A1 2o3 catalysts a~d proposed a mechanism involving methanol decomposition 37.

{2.31) believed to be rate controlling, and the water-gas shift reaction.

{2.51)

The reaction was found to be half order with respect to methanol and water. Disagreement between low carbon monoxide levels and high levels of conversion was attributed to parallel conversion of carbon monoxide to higher alcohols, acetic acid and formaldehyde, though these species were only found in trace amounts. The kinetics of Pour et al. [133] have lately been supported by the findings of Santacesaria and Carra [134] who also used copper-zinc catalysts.

Since the activity of copper for dehydrogenation is higher than its activity for methanol decomposition {sections 2.2.3 and 2.3.2.), a more likely mechanism is that of Kobayashi et al. [135] which involves the dehydrogenation of methanol to methyl formate, decarbonylation to carbon monoxide and methanol, and the water-gas shift. This sequence is represented by

{2.42)

2HCHO + HCOOCH3 (2.45)

(or (2.49)

{2.52) 38.

(2.51)

Takahashi et al. [136] however, showed that the addition of carbon monoxide to the methanol/water feed did not increase the yield of carbon dioxide, contrary to what would be expected by Equation 2.51. The addition of methyl formate, though, did increase the yield. To explain these results, a modified mechanism was proposed involving carbon dioxide formation from formic acid decomposition, the latter being formed by methyl formate hydrolysis.

(1.2)

(2.53)

(2.54)

Formic acid decomposition is known to proceed very rapidly on various metals (including copper) and metal oxides [137,138].

To support their mechanism, Takahashi et al. [136] showed ethyl acetate and acetic acid to be the major products of ethanol steam reforming.

Recent studies of methanol steam reforming have concentrated on the effect of catalyst preparation on the activity of supported copper catalysts. Important insights have been gained into the physical nature of these catalysts. Minochi et al. [139] have shown how pH affects the activity of copper catalysts prepared by alkali precipitation of copper hydroxide. High activity was favoured by 39. high pH [140]. Kobayashi et al. [141] prepared highly dispersed copper-silica catalysts by ion exchange techniques and noted a decrease in turnover frequency and selectivity at high (50 m2/g) copper surface areas. This was attributed to the existence of unreduced copper at high levels of dispersion [141,142].

Inui et al. [143] studied the activity of copper-alumina catalysts for methanol steam reforming. Deactivation was shown to be caused by oxidation of metallic copper during the reaction. Kobayashi et al. [135] prepared copper catalysts mixed with a range of metal oxides and found that Cu-ZnO, Cu-Si02, Cu-cr2o3 and Cu-Sn02 were the most selective at 200°C.

2.3.4. Hydrogenolysis of methyl formate

The reverse of methanol dehydrogenation to methyl formate {equation 1.2) is the hydrogenolysis of methyl formate, given by

{2.55)

For this reason, and because the hydorgenolysis also proceeds over copper catalysts, the reaction will be mentioned here.

Christiansen [144] first studied the hydrogenolysis of methyl formate in the context of the two stage methanol synthesis reviewed in section 2.l.1. Lazier [145] patented a process for methyl formate hydrogenolysis over copper at 180°C, and Brendlein [146] a process using copper chromite catalysts at 1-10 atmospheres and 150-200°C, reporting conversions in excess of 95 percent. High conversions have also been reported by Imyanitov et al. [28] using 40. copper silica and copper-chromium-calcium catalysts at 200-220°C.

Casey [36] has studied the hydrogenolysis of methyl, ethyl, propyl and butyl formates over copper chromite catalysts and, for methyl formate, reported high conversions and selectivity in the temperature range 120-160°C. Higher alkyl formates, for which the hydrogenolysis reaction is:-

HCOOR + 2H 2 + R-OH + CH 30H (2.56) gave lower yields of methanol due to a transesterification reaction

HCOOR + CH 30H + HCOOCH3 + ROH (2.57) between methanol and the formate. Evans et al. [33] have shown how product distribution is largely affected by the rapidly attained equilibrium of Equation 2.57, and have also attributed the diffe­ rences in rates of hydrogenolysis for different alkyl esters to electronic or steric effects. Steady deactivation of copper chro­ mite and Raney copper catalysts for methyl formate hydrogenolysis was attributed to strong chemisorption of methyl formate [32,19].

2.4. Copper Chromite Catalysts

• 2.4.1. Introduction

Copper chromite catalysts were first reported to be active for the hydrogenation of organic compounds by Adkins et al. [147]. Connor et al. [148] and Adkins [149] have outlined the preparation of what will be called 11 Type A11 copper chromite catalysts by the 41.

precipitation of copper-anunonium-chromate. In this method, copper­ ammonium-chromate is precipitated from a solution of cupric nitrate and sodium dichromate by the addition of a11111onium hydroxide accor­ ding to

2Cu(N03)2 + Na 2cr2o7 + 4NH40H J, 2CuNH40HCr04 + 2NaN03 + 2NH4ND3 + H2o (2.58)

The orange precipitate is then calcined at 350°C to yield an equi­ molar mixture of cupric oxide and cupric chromite (Cucr2o4) by the reaction

(2.59)

Barium chromate, claimed to improve catalyst stability, can be co-precipitated from barium nitrate in this method [148,156]. The preparation of a copper chromite catalyst will be described in detail in section 4.2.1.

Another method of preparation in which the copper/chromium ratio can be varied, has been pioneered by Charcosset et al. [150, 151,152] and involves solid state reaction between cupric oxide and chromia at temperatures in excess of 500°C. This reaction is represented by

> 500°C (2.60)

This type of catalyst will be referred to as "Type B11 copper chromite. 42.

Reaction 2.60 is autocatalytic with copper chromate (CuCr04) believed to be an intermediate. Charcosset et al. [152] have proposed the following sequence of reactions to describe the forma­ tion of cupric chromite

(2.61)

(2.62)

(2.63)

Cupric chromite formed by reaction 2.61 is believed to catalyse reaction 2.63 thereby displacing the equilibrium of reaction 2.62 to the right. Banerjee et al. [153] have studied the formation of Type B copper chromites from (a) mixtures of cupric oxide and chromia and (b) co-precipitated gels of copper and chromium hydrox­ ides. A difference in magnetic properties between catalysts prepared by the two methods was noted. The co-precipitation method (b) has been shown by Young et al. to yield active and stable catalysts [154].

The uses of copper chromite catalysts in hydrogenation, dehydrogenation and oxidation have been reviewed by Thomas [155]. Adkins [149] has described the application of copper chromite catalysts to the hydrogenation on a laboratory scale of esters of mono, di and tetra basic acids, esters of aromatic, hydroxy, alkoxy, amino and keto acids, and lactones. The reactions were generally carried out in stainless steel autoclaves at hydrogen pressures of 100-400 atmospheres using methanol or ethanol as solvents. 43.

- Temperatures employed varied in the range 100-300°C depending on the compound being hydrogenated.

A major industrial use of copper chromite catalysts has been in the hydrogenation of vegetable oils in order to improve their storage life, by the hydrogenation of linolenic acid groups, which are known to promote autoxidation [157]. Selective catalysts are needed to hydrogenate linolenic acid to the more stable linoleic acid, an important constituent of "soft" margarines. Catalysts that are too active for hydrogenation, such as nickel, can increase the saturation of the oil, lowering its nutritional value. Koritala and Dutton [158,159] first showed copper catalysts to be highly selective in the hydrogenation of soybean oil at 170-200°C. Copper chromite catalysts were used in a later study of the effect of hydro­ gen pressures (in the range 3-2000 atmospheres) on the selectivity of soybean oil hydrogenation [160]. Okkerse et al. [161] have al so illustrated the superior selectivity of copper over nickel catalysts. Zero-valent copper was deduced to be the active centre, and the induction period observed in the initial stage of hydrogenation was attributed to reduction of copper oxide.

Johansson and Lundin have made a detailed study of soybean oil hydrogenation with copper chromite catalysts [162,163]. The removal of water (formed by reduction of copper) and conjugated dienes by a continuous hydrogen flow was shown to improve catalyst life and activity as both of these species can act as catalyst poisons. Johansson [164] has studied the effect of pressure on the rate of catalyst reduction, and hence the activity for hydrogenation, in soybean oil hydrogenation with copper-chromite. At 185°C, the 44.

maximum rate of hydrogenation was obtained at 6 atmospheres of hydrogen.

The application of copper-chromite catalysts in dehydrogena­ tion processes is not as widespread. Dunbar et al. [165] first reported the activity of copper chromite catalysts for alcohol dehydrogenation, and later showed how catalyst regeneration could be achieved by acetic acid leaching [166]. A supported copper chromite has been patented by Opitz and Urbanski for the dehydro­ genation of c2-c6 alcohols to ketones [167]. The kinetics of ethanol dehydrogenation to acetaldehyde over copper chromite at 225-285°C has been studied by Peloso et al. [87]. de Pinillos and Victor [45] have proposed a process for methanol dehydrogenation to methyl formate using copper chromite catalysts at 200°C.

Recently, the activity of copper chromite catalysts for the oxidation of air pollutants such as carbon monoxide, aldehydes and hydrocarbons has recei.ved some attention [155]. Severino and Laine [168] have shown cupric chromite (Cucr2o4) to be highly active for carbon monoxide oxidation. Heyes et al. [169] have shown how the catalysts can be reactivated for butanal oxidation by washing with water to remove sulphur poisons.

2.4.2. The structure of copper chromite catalysts

The nature of the activity of copper chromite catalysts, and the mechanism by which catalyst reduction occurs, have been the source of considerable contention. Connor et al. [148] first discussed the relationship between activity and structure of the copper chromite catalysts. In the hydrogenation of furfural alcohol 45.

with catalysts prepared via the precipitation of copper anrnonium chromate (Type A copper chromite), reduction of copper oxide to the rnonovalent (Cu(l)) or zero valent (Cu(O)) form was observed, accom­ panied by a decrease in activity. The incorporation of barium, calcium or magnesium chromates during the precipitation was thought to retard catalyst reduction, thereby enhancing activity. Adkins et al. [149,170] believed interaction between cupric oxide and cupric chromite to be crucial to catalyst activity as neither cupric chromite (Cucr2o4), prepared by_ removal of CuO from the catalyst with hydrochloric acid, nor cupric oxide by themselves were active for methyl laurate hydrogenation. In the latter case however, this may have been due to very low active surface area.

The 11 interaction 11 theory of Adkins et al. [170] has been supported by Selwood et al. [171] in studies of the magnetic suscep­ tibility of copper chromite catalysts. Ferromagnetism at low temperatures was exhibited by the catalysts but not by cupric chromite, cupric oxide, nor a mechanical mixture of the two.

Stroupe [171] used X-ray diffraction to study chemical transitions in copper chromite catalysts during reduction and oxida­ tion. At high temperatures(> 900°C) cuprous chromite (cu2cr2o4) was shown to form by two solid state reactions:-

1) (2.64)

2) (2.65) 46.

Type A copper chromite catalysts formed by decomposition of copper­ ammonium-chromate at 350°C (Equation 2.59} were shown by X-ray diffraction to be an intimate mixture of cupric oxide and cupric chromite (designated by 11 Cu0.Cucr20/}. Decomposition at higher temperatures(> 600°C) resulted in the formation of cuprous chromite by reactions 2.64 or 2.65. After usage in liquid phase hydrogena­ tions below 300°C, the catalyst was shown to reduce to metallic copper (Cu(O)) and cuprous chromite (cu2cr2o4). Stroupe attributed the formation of cuprous chromite to reactions 2.64 and 2.65, but could not explain why the excessively high temperatures previously observed in carrying out these reactions were not required for cuprous chromite formation during hydrogenation experiments.

Stroupe proposed a 11 spinel 11 structure for cupric chromite, in agree­ ment with Selwood [171], and further showed how catalyst regeneration could be achieved by oxidation for two hours at 600°C to convert cuprous chromite back to cupric chromite by the reaction

(2.66)

However, catalysts reactivated by this means are not nearly as active as fresh catalysts [149].

Brihta et al. [173] have investigated the claims by Adkins et al. [148,149,170] that deactivation of copper chromite for hydro­ genations was due to reduction of active divalent copper (Cu(II)) to inactive Cu(I) and Cu(O). Raney copper (Cu(O)) was active for hydrogenation of ethyl-oleate and deactivation was thought to be due to reduction to inactive cuprous oxide, cu2o. Thermodynamic 47.

calculations showed that cu2o would not easily be reduced to metal­ lic copper which was suggested to be the true active catalyst. This latter result was in agreement with the proposal of Rabes and Schenck [174] that the active form of Type B copper chromite was finely divided copper supported on chromia and formed by the follow­ ing sequence of reactions:-

(2.67)

(2.68)

This mechanism of formation, however~ is not in agreement with the findings of Charcosset et al. [152] who have shown that reaction between cupric oxide and chromia (Equation 2.60) does not take place below 500°C.

Miya et al. [175] have made a detailed study of the reduction of Type A copper chromite catalysts during the hydrogenation of rapeseed fatty-acid methyl ester at 250 atm and 275°C. Under these conditions, the catalyst was found by X-ray diffraction to reduce to copper on chromia (Cu/Cr2o3) which was shown to have catalytic activity. This observation is in agreement with Rabes and Schenck [174] who proposed that copper on chromia is the active form of the copper chromite catalyst. Water, formed by catalyst reduction during hydrogenation, was strongly adsorbed on chromia, but could be removed by evacuation thereby resulting in a catalyst of superior activity. 48.

The findings of Miya et al. [175] are in conflict with those of Stroupe [172] who identified cuprous chromite (cu2cr2o4) in the reduced copper chromite catalyst. Unfortunately, Stroupe did not specify the hydrogen partial pressure used in catalyst reductions but in light of the experiments carried out at low hydrogen pres­ sures(< 10 atm), it seems likely that low hydrogen pressures were used. Miya et al. [175] have noted that the high hydrogen pressures

(> 100 atm) normally used in liquid phase hydrogenations, are necessary for complete reduction to Cu/Cr2o3.

Reduction of copper chromite catalysts using hydrogen partial pressures of 1 atmosphere and less, and temperatures of 150-200°C, have yielded mixtures of cupric and cuprous chromite, metallic copper and chromia [176,177]. Johansson and Lundin have investi­ gated the mechanism of reduction of copper chromite catalysts during the low pressure (6 atm) hydrogenation of soybean and rapeseed oils [162]. X-ray diffraction, ESCA, and selective extraction of CuO,

CuCr2o4 and cu2cr2o4 were used to monitor changes in catalyst composition over different stages of reduction. Divalent copper was reduced to a mixture of Cu{I) and Cu(O) during the first few minutes of reaction, and Cu(I) continued to be reduced to Cu(O) but at a much lower rate. The following mechanism for catalyst reduc­ tion was put forward:-

(2.69)

(2.70)

(2.68) 49.

Reactions 2.69 and 2.70 were believed to proceed rapidly during the first few minutes of reduction. Reaction 2.68 however, proceeded slowly such that at low hydrogen pressures, the catalytic surface could be considered as Cu{O) supported on cuprous chromite, cu2cr2o4. A subsequent study showed that hydrogen pressures in excess of 20 atmospheres were needed to produce Cu/Cr2o3 by reaction 2.68 [164]. Electron microscopy indicated that while reduction of cupric oxide {Equation 2.69) yielded crystalline copper, amorphous copper appeared to result from reaction 2.68.

Capece et al. [178] have carried out a similar study into the effect of incorporating manganese oxide in the catalyst. X-ray photoelectron spectroscopy {XPS), was used to monitor the change in valency of surface copper during soybean oil hydrogenation at 6 atm and 200°C. Initially only Cu(II) {Cu0.Cucr2o4) was present, but after hydrogenation for 0.5 hours, Cu{I) predominated (Cu2cr2o4). After a further two hours, Cu(O) {Cu/Cr2o3) was the major form of copper present. These findings appear to support the mechanism of Johansson and Lundin [162] described by equations 2.69, 2.70 and 2.68. The effect of manganese oxide appeared to be one of retar­ ding the second stage of reduction (Equation 2.68). Adkins et al. [156,170] have attributed a similar effect to barium oxide when included in Type A copper chromite catalysts, though Boerma has claimed the effect to be due to a higher dispersion of metallic copper [179]. Russian authors have shown barium oxide to form by decomposition of barium chromate during catalyst pre-treatment [180].

Capece et al. [178] also used XPS to study the fate of surface chromium during catalyst pre-treatment, though the experi­ ments were hindered by agglomeration of copper on the surface. 50.

Although Cr3+ should predominate in Cucr2o4, Schreifels et al. [181] have shown by XPS that surface ratios of cr6+/cr3+ as high as 2/1 can exist for fresh copper chromite catalysts. d'Huysser et al. [182] have identified the existence of cation vacancies associ­ ated with Cr6+ ions which are localized at octahedral sites in the cupric chromite spinel. Schreifels et al. [181] have also shown transition of cr6+ to cr3+ by calcination in air at 500°C, to be retarded by the presence of barium oxide, BaO.

Reduction of Cu0/Cr2o3 mixtures, prepared by the co-precipi­ tation method of Young et al. [154], has been studied by Banerjee and Naidu [183,184] using thermal analysis. Reduction of cupric oxide was shown to convnence at temperatures around 150°C. The formation of cuprous chromite, observed at 500°C, was attributed to the reactions:-

(2.71)

1 2Cr0 + 2 o2 (2.72)

Chromia was shown to be important in maintaining high surface areas and a 1:1 ratio of CuO to Cr2o3 was found to be very suitable for nitrobenzene, hydrogenation.

Choudhary and Sansare (185] have described a gas chromato­ graphic technique for studying hydrogen adsorption/desorption from copper chromite catalysts. Reversible and irreversible hydrogen adsorption was observed over the temperature range 30-350°C, and three types of active sites, copper, chromia and copper-chromia 51.

co-sites, have been postulated [186,187]. Aissi et al. [188] have reported the existence of reactive hydrogen species in the copper chromite spinel using thermogravimetry and wide line NMR. 52.

CHAPTER 3

PROJECT OBJECTIVES

The objectives of this study were:-

(i) To investigate the structure of the 11 copper chromite 11 type of catalysts, with particular referenc~ to the nature of the active surface;

(ii) To establish reaction conditions whereby the activity of copper catalysts for the dehydrogenation of methanol to methyl formate can be investigated;

(iii) To compare the activity of various catalysts for the reaction, and to account for differences in activity, selectivity and stability;

(iv) To investigate the mechanism of the dehydrogenation reaction on copper. 53.

CHAPTER 4

EXPERIMENTAL TECHNIQUES

4.1. Catalyst Testing

4.1.1. Apparatus

The flow apparatus built for catalyst testing is schemati­ cally illustrated in Figure 4.1. Catalyst particles (0.5-2 g} were loaded into a stainless steel U-tube reactor (e) (400 mm x 9 nm O.D.) which was ifllTlersed in a salt bath (f) containing a molten eutectic of potassium nitrate (53 weight percent), sodium nitrite (40 weight percent) and sodium nitrate (7 weight percent). To facilitate heat transfer, the reactor was packed with 1-2.5 mm glass beads.

The salt bath was stirred with an air motor driven by compressed air (400-500 kPa), and heated using a 1500 kW coiled electrical heating element. Temperature was controlled to± 0.5°C with a Shinko proportional-integral controller via a stainless steel-sheathed, chromel-alumel thermocouple which was irrmersed in the salt bath. A second thermocouple (also chromel-alumel) connected to a digital voltmeter, was used to measure the salt bath temperature to± 0.25°C. Calibration was carried out with an ice-water bath and 0-360°C mercury thermometer. Experiments conducted with thermocouples inmersed in the catalyst bed showed that the temperature differential between the salt bath and the catalyst bed never exceeds 1.5°C. 54 .

. u ~'4------f______

a,

u

a, :r: z.. .0

N I

0

Figure 4.1. Experimental apparatus for catalyst testing: (a) Pump; (b) Rotameters; (c) Pressure Gauge; (d) Pre-heater; (e) Reactor; (f) Salt Bath; (g) Back-pressure regulator. 55.

Methanol was pumped to the reactor using an Eldex Model E-120-S-2 pump {a) at flowrates in the range 0.3 and 4 cm3min- 1. A needle valve, connected on the outlet of the pump, was used to establish a back-pressure of 1000-2000 kPa {read from a 0-25000 kPa Bourdon gauge) as required for efficient pump operation. The reservoir for the pump was a 100 cm3 burette graduated to 0.1 cm3 divisions, and accurate measurement of the methanol flow rate {± 2 percent) was achieved by timing the fall of the methanol meniscus {generally over 5-10 minutes) using a stopwatch. When methyl formate, or methanol isotopes had to be pumped, syringe pumps of 2 cm3 capacity, were generally used. These pumps, built in the laboratory, utilized a Phillips synchronous motor to drive the teflon plunger of a precision bore syringe. Liquid feed rates were varied by the substitution of different Phillips chart­ recorder gear boxes. The syringe barrels were graduated to 0.001 cm3 and flowrates {0.05, 0.1 and 0.2 cm3min- 1) were accurate to± 2 percent.

Gases {hydrogen, nitrogen, helium, carbon monoxide or deuterium) could also be introduced to the reactor. Flowrates (20-1000 cm3min- 1) were controlled by needle valves and measured with rotameters which had been calibrated using a wet gas meter and stopwatch. Control of gas flowrate was accurate to± 5 percent. The two gas lines were combined at the inlet to a Fairchild 0-200 kPa back pressure regulator {g) {maintained at 50 kPa gauge) to ensure that flowrates did not vary with pressure-drop across the reactor. 56.

Liquids were vaporized and mixed with gases in a preheater {d) consisting of a stainless steel tube (180 nm x 40 nm O.D.) which was inmersed in the salt bath and located inmediately upstream of the reactor. A Bourdon gauge (c) (0-100 kPa gauge) was used to measure the pressure drop across the reactor. When the syringe pumps were used, the pre-heater was considered unnecessary due to the low flowrates involved (300 cm3min-l maximum) and was bypassed to minimize the system dead volume. In this case reactants were fed directly to the inlet side of the reactor in order to reduce the holdup of expensive isotopically labelled chemicals.

In order to prevent condensation of methanol/methyl formate vapours, the gas lines between the preheater and the reactor, and between the reactor and the sample valve, were heated to 80-100°C with teflon coated wire connected to a variac. The sample valve (Valeo 6 port, teflon stem) was mounted in an insulated box main­ tained at 150°C with a 75 watt light globe. Sample loops of 0.37 cm3 or 1.63 cm3 were used to inject a sample of gas from the reactor into the hydrogen carrier gas of a Gow Mac series 550 Gas Chromatograph fitted with Thermal Conductivity detector. The conditions of gas chromatographic analysis will be discussed in section 4.1.2.

After leaving the valve box, gases were passed through a Dreschel bottle, where liquids were condensed, and then vented. For more quantitative collection, as required during isotope experiments, gases were passed through a pyrex tube inrnersed in an isopropanol-dry ice bath (- 78°C) and connected to the gas lines by means of a compression fitting. A gas sample could be taken in 57. a 1 m x 6.25 mm O.D. copper tube fitted with on-off valves at either end, and connected downstream from the pyrex tube.

Fonnaldehyde was fed to the reactor by connecting a gas line from the back-pressure regulator (g) to the inlet of a U-tube filled with paraformaldehyde and immersed in an oil bath at temperatures in the range 80-120°C. This U-tube was then connected to the preheater via lines that were electrically heated to 150°C. The flowrate of formaldehyde was controlled by varying the temperature of the oil bath thereby detennining the partial pressure of fonnaldehyde above the solid paraformaldehyde. The properties of fonnaldehyde poly­ mers, and the decomposition or 11 depolymerisation 11 into gaseous formaldehyde, have been examined by Walker [190] and Melia [191]. Generally, the oil bath was set at 110°C for 29 minutes to remove water and methanol from the polymer, and then cooled to 80°C in order to obtain a formaldehyde partial pressure of 0.3 atmospheres. This was in agreement with the vapour pressure data of Nielsen and Ebers [192].

All lines in Figure 4.1 were of 3.125 mm O.D. stainless steel tubing except for the line between the pump (a) and the preheater

(d) which was of 0.0625 mm O.D. stainless steel tubing. 11 Swagelock 11 connections were employed.

Sources and purity of the reagents described in this section are shown in Table 4.3.

4.1.2. Chromatographic Analysis

Carrier gas flowrates were maintained constant in the range 30-50 cm3min-l with a Porter mass-flow controller. Different 58.

flowrates were used depending on the column employed and the range of products being analysed. Conditions of analysis for the two columns, along with typical retention times for the reaction pro­ ducts, are shown in Table 4.1.

Peak areas were measured from the chromatographic response using a Hewlett Packard 3390A integrator. Molar responses (relative to methanol) were calculated for carbon monoxide, carbon dioxide, water, formaldehyde and methyl formate by injection of known volumes of these compounds into the chromatograph. The values obtained agreed closely with those of Messner et al. [189] (Table 4.2)

TABLE 4.1. Retention times for gas chromatographic analysis

Column Carrier Gas Temperature Component Retention Flowrate Time (cm3min-1) (oc) (minutes)

3.6 m X 40 130 Carbon Monoxide 0.69 3.2 nm O.D. Carbon Dioxide 0.84 Porapak-Q Water 1.27 Methanol 1.78 Methyl Formate 3.12

1 m X 50 130 Carbon Monoxide 0.31 6.3 nm O.D. Carbon Dioxide 0.46 Porapak-N Formaldehyde 1.18 Dimethyl Ether 1.79 Water 2.22 Methanol 2.80 Methyl Formate 4.59 59.

TABLE 4.2. Relative molar responses {to methanol)

Compound Experimental Messner et at. [189]

Methanol 1.0 1.0 Carbon Monoxide 0.80 0.77 Carbon Dioxide 0.87 0.87 Water 0.50 0.39 Formaldehyde 0.78 Methyl Formate 1.20

4.1.3. Sources and purity of chemicals employed

TABLE 4.3.

Material Source Purity {%)

Hydrogen C. I.G. Ltd. 99.995 (Chromatograph) 99.9 {Reactor) Nitrogen C. I.G. Ltd. 99.9 Helium C. I.G. Ltd. 99.9 Carbon Monoxide Matheson 99.9 Methanol N.S.W. Government Stores 99.8 Methyl Formate Ajax Chemi ea l s 98 Parafonnaldehyde Foseco 91 Isotopes CH30D Merck 99 CD30H Merck 99 CD30D Merck 99.6 Deuterium Matheson 99.5 60.

4.1.4. Calculations

Comparisons between catalysts are usually made on the basis of conversion, initial reaction rate and selectivity under a standard set of conditions where one or more variables (such as temperature, flowrate or pressure) are held constant. For the majority of tests outlined in this thesis, catalyst comparisons are made at constant temperature, and conversions and selectivities are obtained for a range of methanol flowrates.

The dehydrogenation of methanol over copper produces carbon monoxide and carbon dioxide as well as methyl formate, and calcula­ tion of conversion and selectivity depends on knowledge of the chemical reactions involved in producing these components. If no byproducts are produced, then the stoichiometry of Equation 1.2 gives:-

X = (4.1) YMeOH + 2xYMeF where "X" is conversion of methanol and YMeOH, YMeF are mole fractions of methanol and methyl fonnate respectively, in the product stream of the reactor. Mole fractions are calculated for carbon containing species in the reactor effluent. In this study, carbon dioxide is presumed to be formed from trace water via the sequence of reactions outlined in section 2.3.3. This sequence is given by

HCOOCH3 + H2o ~ HCOOH + CH30H (2.53)

HCOOH ~ co2 + H2 (2.54) 61.

Carbon monoxide is believed to form via the decarbonylation of methyl formate outlined by Christiansen [144], Schwab and Knoezinger [193] and Higdon et al. [44] and occurs by the reaction

(2.12)

Formation of carbon monoxide by methanol decomposition is not likely to be significant below 300°C (Section 2.2.3).

To account for the methyl formate that has further reacted to carbon monoxide and carbon dioxide, Equation 4.1 is rewritten as:

Yeo+ Yco2 + 2xYMeF x-----,~--- (4.2) - YMeOH + 2xYMef where YCO' Yeo are mole fractions of carbon monoxide and carbon 2 dioxide respectively. The extra molecule (of methanol) produced in Equations 2.53 and 2.12 is incorporated in the observed value of YMeOH. Selectivity, S, is now defined by dividing the apparent conversion to methyl formate (Equation 4.1) by the true conversions (Equation 4.2), or

2xYMeF S=~-~----=-~- (4.3) Yco + Yeo + 2xYMeF 2

The values for mole fractions required for calculation of conver­ sion and selectivity are obtained by integration of the chromato­ graph response and division by the appropriate relative molar response, as outlined in the previous section. These are not the 62. true mole fractions of course, as hydrogen is not detected, but as conversion and selectivity are based on a carbon balance, the presence of hydrogen is irrelevant. The true mole fractions are calculated by knowing the total number of moles (NT) at a given conversion, which, based on Equation 1.2, is given by

(1.2)

NT= 2-2X + X + 2X = 2+X (4.4)

The total moles of carbon containing species, Ne is then given by

N = (2+X) - 2X = 2-X C (4.5)

The true mole fractions (Y) can now be obtained as

I Y = Y X (2-X)/(2+X) (4.6)

I where Y is the apparent mole fraction based on carbon containing species only.

For the majority of .catalysts tested, conversion and selec­ tivity are plotted against flowrate of methanol expressed as a liquid Hourly Space Velocity (LHSV) where:

LHSV (h-1): Flowrate of methanol h-l Mass of catalyst g (4.7) 63.

This definition enables convenient comparison between catalysts of widely differing bulk densities.

Although the effect of pressure is not considered in this study, total pressures have been calculated for all catalyst compa- I risons. The pressure drop (P) across the catalyst bed was assumed to be linear and an average pressure was given by:

I P = 1 atmosphere+ P /2 (4.8)

A maximum value for P of 1.3 atmospheres was encountered (equivalent to a pressure drop of 0.3 atmospheres).

Comparison between catalysts was also made on the basis of reaction rate. For low conversions (< 10 percent) the characte­ ristics of a differential reactor were assumed and rate was given by:-

r = FMeOH X w (4.9) where r = rate of methanol consumption (mole h-lg-1)

F = flowrate of methanol (mole h-1)

X = conversion of methanol

w = mass of catalyst (g)

Calculation of rate was achieved by plotting conversion (x) against the inverse of flowrate (F- 1) and, as shown by Figure 4.2, a straight line is obtained at low conversions(< 7 percent) indica­ ting the validity of the initial rate approach. The slope of this line yields the reaction rate. It will be shown that zero order 64.

25

(!) 20

15

,-,. :-....: ...... , 10 z ,_.~ c.n 0::: w > z 5 ~ u

0

0 20 40 60 80 100 1/F CMIN/MCJLEJ

FIGURE 4.2. Graphical determination of initial rate 65. kinetics are obeyed and hence rate and rate constant are equivalent. Specific rate constants then can be calculated as follows:-

Specific rate constant (mole h-lm-2)

_ rate (mole h-lg-1) (4.10) - metal surface area (m2g-1)

The measurement of metal surface area will be discussed in Section 4.3.2.2.

The calculation of conversion, ~electivity, rate, pressure, mole fractions and LHSV from peak areas and methanol flowrate are incorporated in a BASIC computer program listed in Appendix III.

4.2. Catalysts

4.2.1. Copper chromite catalysts

The three colllllercial copper chromites that were used provided an adequate representation of the types of catalysts discussed in Section 2.4. Suppliers and nominal compositions are listed in Table 4.4. For two of the catalysts, the mass fractions of the catalyst components do not add up to 100 percent. This is presumably due to binder and other components used in the prepara­ tion. Peloso et al. [87] have thoroughly analysed a Harshaw copper chromite catalyst and found it to contain 9.3 percent Si02, 3.3 percent Na2o and 12.8 percent binder, as well as 41.2 percent CuO and 33.4 percent cr2o3• A more detailed elemental analysis of the catalysts is given in Chapter 5. The catalysts were supplied as 6x4 mm pellets and were crushed and sieved to 400-600 micron particles before use. 66.

TABLE 4.4. Suppliers and nominal compositions of copper chromite catalysts

Catalyst Source Nominal Composition {wt%)

1808 Harshaw 43% CuO; 38% Cr2o3 0203 Harshaw 80% CuO; 20% Cr2o3 G-22 Girdler 31% Cu; 25% Cr; 10% Ba

A Type A copper chromite catalyst was prepared as a standard by the method of Miya et al. [175]. Sixty cm3 of a 28 percent ammonia solution was added to a solution of 60 g of Na 2cr2o7.2H2D in 260 cm3 of distilled water. Over a period of 15 minutes, a solution was added containing 100 g of Cuso4.5H2o in 300 cm3 of distilled water. Temperature was maintained at 80°C and vigorous stirring applied for a period of 1.5 hours. The orange precipi­ tate, formed by the reaction:-

2CuS04 + Na 2cr2o7 + 4NH3 + 3H20 ! {4.11) was filtered and washed thoroughly with distilled water before being dried at 110°C for 16 hours. The powder was then calcined at 400°C for 30 minutes and cupric oxide and cupric chromite were formed by reaction 2.59.

Pure cupric chromite, Cucr2o4, was obtained by leaching this catalyst according to the method of Adkins et al. [170]. Teng of 67. the catalyst was added to 75 cm3 of 32 percent hydrochloric acid and heated to just below boiling for one hour. The residue was separated by filtration and washed with 10 percent acetic acid before being dried at 110°C for 16 hours. As will be shown in Chapter 5, the X-ray diffraction pattern and elemental analysis of the residue corresponded to Cucr2o4. The filtrate was retained for determination of cupric oxide content. The three commercial cata­ lysts were also treated by this method.

4.2.2. Supported catalysts

Copper catalysts supported on silica, alumina, chromia and magnesia were prepared by impregnation and ion-exchange techniques. The latter technique and the chromia and magnesia supports, were developed in this study, and will be described in Chapter 7. The impregnation method can be generalized for all supports and so will be described here. Generally 4 x 6 mm pellets were supplied and crushed and sieved to 400-600 micron particles prior to use. Before impregnation, support particles were heated at 110°C for 16 hours and then evacuated for one hour to remove moisture. Ten­ twenty g of support was impregnated at one time. Impregnation was perfonned by the incipient wetness technique described by Emmett [194]. Water was added dropwise from a burette to a known mass of catalyst, with thorough mixing. The volume of water required to "wet" the support (i.e. for the particles to agglomerate together) was recorded. Generally, a high surface area support required 1 to 1.2 ml of water per gram of catalyst for "wetting". Once the wetting volume was known, a volume of cupric nitrate stock solution (634 g 96 percent Cu(N03)2.3H20 per 1000 cm3) containing the 68.

required mass of copper could be taken and diluted to a volume of solution that would uniformly fill the catalyst pores. This solution was then added dropwise to the support, with thorough mixing, and the wetted support dried at 110°C for 16 hours before calcination at 350°C for three hours. A typical set of calcula­ tions involved in the preparation of 10 g of 5 percent copper on silica catalyst, are shown in Appendix IV.

Characterisation of supported catalysts is carried out in Chapter 7. Most of the supported catalysts studied were prepared using a silica support (Davison Chemical Division of W.R. Grace and Co. - Code ID 57), with a nominal surface area of 300 m2g-l and pore volume of 1 cm3g-1, although a low surface area silica support (50 m2g-1) and an alumina support (200 m2g-1) were also used.

4.2.3. Raney copper catalysts

Copper-aluminium alloys used in the preparation of Raney copper catalysts were obtained as a comnercial alloy from Ingot Metals Pty. Ltd. or were prepared by the method of Marsden et al. [196] by adding aluminium prills to molten copper in the carbon crucible of an induction furnace. After stirring with a carbon rod for 15-20 minutes, the molten alloy was quenched in water. Cu-Al-Zn and Cu-Al-Cr alloys were also prepared in the laboratory by a similar method. All alloys were crushed in a jaw crusher and sieved to 400-600 micron particles.

Catalysts were prepared from alloy particles by leaching with caustic soda to remove aluminium, by the method of 69.

Marsden et al. [196]. The leaching process proceeds by the chemical reaction [197]

- - 3 Al + OH + 3H 20 4 Al(OH) 4 + 2 H2 (4.12)

Alloy particles (20 g) were placed in a conical flask containing 100 cm3 of water and irrmersed in a 50°C water bath. After tempe­ rature equilibration was achieved, 40 percent weight/weight caustic soda was added according to the schedule in Table 4.5. As the reaction is highly exothermic, the initial rate of addition is low

to prevent the temperature of the mixture exceeding 50 ± 3°C. After three hours at 50°C, at which time the reaction is virtually complete (as shown by measuring hydrogen evolution), the spent caustic is decanted, replaced with 100 ml of fresh 40 weight per­ cent caustic solution and left for a further three hours to remove residual high surface area alumina. The solution is now decanted

TABLE 4.5. Schedule for caustic soda addition in the preparation of Raney catalysts

Rate of Time Volume of NaOH addition added (ml/min) (minutes) (mls)

0.2 10 2 0.5 20 5 1.0 25 5 2.0 29 8 Balance at 5 mls/min 41 81

-....J -....J

0 0

/g) /g)

1.2 1.2

2.6 2.6

3.9 3.9

2

8.6 8.6

18.1 18.1

31 31

(m

Copper Copper

Areas Areas

19.9 19.9

15 15

21.2 21.2 12.0 12.0

18.6 18.6

45 45

• •

Total Total

Surface Surface

Properties Properties

Catalyst Catalyst

Content Content

79.6 79.6

63.8 63.8

95.6 95.6

99.3 99.3

(wt%) (wt%)

Copper Copper

3 3

3 3

3 3

0.5 0.5

3 3

62. 62.

(h) (h)

Time Time

Leaching Leaching

catalysts. catalysts.

10 10

40 40

40 40

40 40

40 40

40 40

NaOH NaOH

(wt%) (wt%)

Ltd.) Ltd.)

Concentration Concentration

copper copper

Pty. Pty.

Raney Raney

10 10

Cr Cr

of of

hours hours

Catoleum Catoleum

3 3

by by

15 15

Zn Zn

after after

Characteristics Characteristics

NaOH NaOH

Composition Composition

(Supplied (Supplied

50 50

(wt%) (wt%) 49 49

48 48

Al Al

2 2

6. 6.

4. 4.

Fresh Fresh

Alloy Alloy

CuA1

2. 2.

1. 1.

41 41

35 35

50 50 50

50 50 50

50 50 50

Cu Cu

TABLE TABLE ~2 ~2 71.

and the catalyst washed with distilled water until a wash-water pH of 7 is obtained. Catalysts are stored under water at 3°C.

As can be seen from Table 4.6, where alloy compositions, leaching conditions and catalyst characteristics are listed, a high surface area catalyst of almost pure copper is readily obtained, and it was this type of catalyst that was most often used. However, experiments were also performed with the chromium and zinc contain­ ing catalysts, and catalysts of varying copper surface areas obtained by altering leaching times and caustic soda concentration.

The determination of surface areas and copper content are outlined in section 4.3.

4.2.4. Other catalysts

Miscellaneous catalysts tested, along with suppliers and composition, are listed in Table 4.7. One gram of 400-600 micron particles was generally tested.

TABLE 4.7. Miscellaneous catalysts

Catalyst Supplier Nominal Composition (wt%)

Cu0/Zn0/Al 2o3 United 71% Cu; 22% ZnO; 6% Al 2o3 (C79-4) Catalysts Zinc Chromite Vulcan 25.8% ZnO; 74.2% ZnCr2o4 ZC-105 Cincinnati Cu/Ni0/Al 203 Phillips 9.5% Cu; 0.5% NiO on Al 203 Pd/Al 2o3 ICI Aust. 1% Pd on Al 2o3 Ag0/Al 2o3 Harshaw 11% on Al 2o3 72.

4.2.5. Catalyst pre-treatment

Copper oxide was inactive for methanol dehydrogenation, so all catalysts were reduced in pure hydrogen at 220°C prior to activity testing. A minimum reduction period of three hours could be used at a hydrogen flowrate of 100 cm3min- 1• Generally however, reduction was carried out overnight (16 hours). Voge and Atkins [195] have studied the rate of reduction of supported copper cata­ lysts with hydrogen and identified an induction period due to initial growth of Cu(O) nuclei. Reduction is then promoted by hydrogen adsorption on these nuclei and reaction occurs at the Cu/CuO interface.

Raney copper catalysts, which do not need to be reduced, were loaded wet into the reactor. To remove water, a flow of hydrogen (60 cm3min- 1) was passed through the reactor which was then inrnersed in an oil bath at 70°C. The temperature of the bath was then raised to 110°C over a period of one hour, after which time the reactor was capped and transferred to the salt bath.

4.3. Catalyst Characterisation

4.3.1. Atomic absorption analysis

Copper loadings of all catalysts were determined by Atomic Absorption Spectroscopy (AAS) using a Varian-Tecktron AAS Spectrophotometer.

In the case of supported copper, and Raney catalysts and alloys, approximately 0.5 g of sample was dissolved in 100 ml of a 3:1 mixture of concentrated hydrochloric and nitric acids. All 73. of the copper was dissolved after five hours. The cupric oxide content of copper chromite catalysts was determined by analysis of the filtrate obtained in the leaching procedure of section 4.2.1. For the determination of total copper in these catalysts, a diges­ tion procedure developed in the School of Geology, University of New South Wales, was employed. In this procedure, approximately 0.25 g of sample is weighed into a teflon beaker to which is then added 25 cm3 of 40 percent hydrofluoric acid and 10 cm3 of a 2:1 mixture of nitric and perchloric acids. After leaving at 95°C overnight, the sample is heated to dryness. Care is taken to avoid baking the sample. Ten cm3 of 50 percent hydrochloric acid is then added, and after heating for 30 minutes, the mixture is washed into a volumetric flask and made up to volume with deionized water.

Standards (2, 4, 6 and 8 ppm) were prepared from cupric chloride, and solutions of dissolved catalyst were diluted to be within the concentration limits of the standards. Concentrations were then determined by fitting absorbance readings to a calibra­ tion line. Conditions for AAS analysis of copper in the range 0-10 ppm are:-

Slit Width - 0.2 nm Flame - air-acetylene oxidizing Wavelength - 324.7 nm Lamp Current - 3mA

4.3.2. Surface area characterisation

4.3.2.1. Measurement of total area by nitrogen adsorption

As catalytic activity can depend on the abundance and 74.

availability of active surface, the measurement of surface area and pore size is of considerable importance. The significance of pore structure and surface area in heterogeneous catalysis has been reviewed by Enmett [198], Kramer [199] and Wheeler [200]. Innes [201] has examined the techniques available for the study of these catalyst properties.

In this study, surface area was determined by measuring the quantity of nitrogen adsorbed onto the catalyst surface at succes­ sively higher pressures while maintaining temperature constant at -196° (an adsorption isotherm). The data was fitted to the rela­ tionship of Brunaur, Enmett and Teller, which in linearized form is given by:

P 1 C-1 P (4.13) V(Po-P) = vmc + vmc + Po where P is the equilibrium pressure of nitrogen, Po is the saturation pressure of nitrogen at -196°C and C is a constant. Both C, and the volume of the nitrogen monolayer, Vm, can be obtained from the slope and intercept of Equation 4.13, and the surface area is obtained by multiplying Vm by the cross-sectional area of a nitrogen molecule (15.4 A2 molecule-1). The BET equation 0 (4.13) is valid for pore widths between 20 and 500 A

By measurement of the desorption of nitrogen at successively lower equilibrium pressures (a desorption isotherm), the pore size distribution can be obtained using the Kelvin equation which is given by [203]: 75.

r _ 4.14 - log (Po/P) (4.14)

where r is the pore radius (cylindrical pores assumed). In this study, pore size distribution is presented as a plot of cumulative surface area against pore radius in order to show how much of the surface area is related to a particular range of pore sizes. Cumulative surface area is calculated from P and Po by the method of Gregg and Sing [203].

The determination of adsorption and desorption isothenns was carried out by the classical fixed volume technique outlined by Innes [201]. A Micromeritics 2100E ORR surface area/pore volume analyser, or Micromeritics Accusorb instrument was used. Rapid surface area measurements were obtained on a Micromeritics High Speed Surface Area Analyser using nitrogen or argon as an adsor­ bant. This constant pressure technique measures adsorption at a value of P/Po high enough to ensure monolayer coverage but low enough to avoid multilayer adsorption or capillary condensation. A value of P/Po= 0.2 is general1y employed.

All catalysts were dried at 110°C for 16 hours prior to surface area determination, and in the case of the fixed volume technique, evacuated for the same period. Copper chromite catalysts were pre-reduced for the fixed volume technique. A Fixed-Pressure Flow technique for total surface area detennina­ tion by nitrogen adsorption and metal surface area by reaction with nitrous oxide that enables in situ reduction to be carried out, will be outlined in the following section. 76.

4.3.2.2. Measurement of copper area by reaction with nitrous oxide

Detennination of metal surface area usually involves chemi­ sorption of a gas selectively onto the metallic component of a catalyst, with metal surface area being calculated from the amount of gas adsorbed and the cross-sectional area of the chemisorbed species. Scholten [204] has reviewed the numerous chemisorption techniques available for the study of metal surface areas.

Carbon monoxide, hydrogen, oxygen and nitrous oxide are known to chemisorb onto copper surfaces. In the case of carbon monoxide however, differentiation between physical and chemical adsorption is difficult due to the low heat of chemisorption [204, 205]. Hydrogen chemisorption has also been employed [206,207] but results are unreliable due to the very low rates of hydrogen uptake [205]. Oxygen chemisorbs rapidly onto copper but the high heat of reaction promotes bulk oxidation [208].

Dell et al. [208] first studied the reaction of nitrous oxide on copper and showed the reaction to proceed as

(4.15)

Osinga et al. [209] first proposed a static method for copper surface area determination by nitrous oxide chemisorption at 298 K. Scholten and KoYaJinka [210] showed fractio~al surface coverage with nitrous oxide to increase from 0°C to 100°C and that bulk oxidation began at 120°C. Best results were achieved at 90°C though Dvorak and Pasek [211], and Sengupta et al. [212] have proposed that bulk oxidation can begin at 70°C. 77.

Measurement of copper surface area by the method of Scholten and Kovlinka [210] has been successfully applied by Iglesia and Boudart [213] to the study of copper catalysts for fonnic acid decomposition (Equation 2.54), and by Takezawa et al. [140,141,142, 214] to the study of copper catalysts for steam refonning of methanol (Equation 2.50). In a related study, Takezawa et al. (215] have used UPS to verify the stoichiometry of reaction 4.15 and have shown bulk oxidation to begin at 100°c.

Recently, Evans et al. [32,216] have used a chromatographic technique for the determination of copper surface areas of a wide range of copper based catalysts. For copper powder and Raney copper, an accurate estimation(± 5 percent) of copper surface area could be obtained at 90°C using a single pulse of excess nitrous oxide. When catalysts containing zinc and chromium oxides were used, bulk oxidation was enhanced. In the case of copper chromite however, the use of lower temperatures resulted in incomplete surface coverage. It was felt that the most accurate results could be obtained by using one pulse containing a 3-4 times excess of nitrous oxide at 90°C. Under these conditions, complete surface coverage could rapidly be obtained and over­ estimation of copper surface area due to the slow bulk oxidation process could be kept well below 10 percent.

Experimental

The experimental apparatus shown in Figure 4.3 is the same as that employed by Evans et al. [32,213]. A 6-port Valeo sample valve {l) was used to inject a 1-3 cm3 sample of nitrous oxide co

-....J

Sanple

H2

-

valve;

,.

area:

sample

column

valve.

gas surface

He

Valve

loop

copper

,

switching

Porapak

dual

of

-

3-Way

6-Port

®-1•

Vent

10-Port,

3

3

4.

measurement

2.

(e)

t•>---a@,

I

the

•I

cell;

I

for

valve;

•1

sample

apparatus

conductivity

gas

6-Port

Thermal

1.

Experimental

3.

4.3.

Figure 79.

(C.I.G. 99 percent) into a helium carrier gas (C.I.G. 99.999 per­ cent) maintained at a flowrate of 30 cm3min- 1• The pulse then flowed through one side of a thermal conductivity detector (3) consisting of four thermistors connected in a Wheatstone bridge. Incoming carrier gas determined the baseline responses of the detec­ tor. The pulse of nitrous oxide then passed into a chromatograph

oven (90 ± 1°C) and over 0.1-0.3 g of catalyst contained in the void space of Swagelock 1/4 inch to 1/8 inch reduced fitting. A 1 m x 3.125 mm column of Porapak T maintained at 25°C was used to separate the pulse into nitrous oxide and nitrogen before it passed back through the detector and to vent. The detector response was connected to a Spectra Physics Minigrator which was used to measure the areas of the peaks due to the inlet nitrous oxide and the peaks due to nitrogen and nitrous oxide in the product mixture. The 6-port Valeo valve (4) was used to carry out in situ reduction by directing hydrogen over the catalyst sample.

Calculations

Calculation of copper surface area first involved subtracting nitrogen impurity in the nitrous oxide, from the outlet nitrogen peak by

N2 = N21 - RxN20(I) (4.16) where N2 = nitrogen area N 1 2 = observed nitrogen area N20(I) = inlet nitrous oxide area R = mole fraction of nitrogen impurity(= 2.54xlo-3). 80.

The volume of nirous oxide (VN 0) that has reacted is given by 2

(4.17)

where L = sample loop volume (cm3)

Copper surface area (Seu) is finally given by

VN 0 = 2 (4.18) WX 0 where W = mass of catalyst (g)

D = site density of copper (0.298 cm3m- 2 [32,216]}.

4.3.2.3. Single-point nitrogen adsorption

Single-point fixed-pressure surface area determination similar to that described by Nelson and Eggertsen [217], was carried out on the apparatus shown in Figure 4.3 by employing a 30 percent mixture of nitrogen in helium as carrier gas. Monolayer coverage is achieved at this partial pressure of nitrogen [218]. The 10-port Valeo valve {2) was used to inject two different but accurately known pulses of nitrogen (range 1-3 cm3) into the carrier gas (30 cm3min-1) and a calibration factor (peak area per cm3 of nitrogen) could be obtained from the detector response. After the pre-reduced sample had cooled to room temperature, the sample tube was immersed in a Dewar flask containing liquid nitro­ gen and a broad negative peak was obtained due to the adsorption of nitrogen onto the sample. When the detector response had stabilized, the sample tube was inrnersed in hot water and a 81.

detector response corresponding to a sharper peak of desorbed nitrogen was observed after one minute.

The surface area (S) was calculated from the relationship [215]

2.79 x VN 2 s = w (4.19)

where VN = volume of nitrogen desorbed (cm3) 2 determined from the calibration factor and

W= mass of catalyst (g).

Agreement between surface areas obtained by this technique, and by the techniques outlined in section 4.3.2.1., was better than 5 percent.

4.3.3. X-ray Diffraction

X-ray diffraction (XRD) has been used extensively in scien­ tific research to relate the physical properties of solids to the arrangement and spacing of atoms within crystals. It is also a practical tool for the qualitative and quantitative identification of crystalline compounds. Cullity [219], and Klug and Alexander [220], have extensively reviewed the theory and application of X-ray diffraction.

The primary use of XRD in this study is in the identifica­ tion of solid phases within the copper chromite catalysts. Similar work has been done by Stroupe [172], and Johansson and 82.

Lundin [162], but under markedly different conditions of catalyst pre-treatment and use. Identification of crystalline species is normally carried out by comparing the XRD spectra with the exten­ sive files of major diffraction peaks for inorganic compounds [221]. In the case of cupric and cuprous chromites however, the major peaks are not dissimilar to those for cupric and cuprous oxide, and the full spectrum, for use as a standard was required. Complete spectra for cupric and cuprous chromite have been presented by Schulz et al. [222], however d spacings were not given and considerable error is involved in the estimation of d spacings from the figures presented. To assist in the interpretation of this data, a cupric chromite standard was prepared by the method outlined in Section 4.2.1. XRD has been used for qualitative analysis of other types of copper catalysts by Friedman et al. [223] and Lo Jacono et al. [224] in the study of copper on alumina, and by Shimokawabe et al. [142,225] in the study of copper on silica.

XRD was also used to measure copper crystallite sizes for the majority of copper catalysts studied in this thesis. The method involves the use of the Scherrer equation which relates peak broadening to crystallite thickness and is given by [219]:

L = 0.9 ). B Cose (4.20)

0 where L = crystallite size (A)

0 ). = X-ray wavelength (A)

B = peak width at half maximum intensity for the diffraction peak (2e) 83.

0 = Bragg angle (in radians) for that peak.

To allow for the breadth of the diffraction line, a standard with

0 crystallite size greater than 1000 A is added to the sample. The width of the diffraction line for the standard approximates that due to the instrument and a correction is made using the Warren relationship which is given by [226]:

82 = B 2 - B 2 M s (4.21) where B = true peak width at half maximum intensity

BM = the observed value

Bs = that of the standard.

Iglesia and Boudart [213] have recently used XRD for crystallite size determination in copper catalysts.

A Rigaku Geiger Flex Diffractometer utilizing CuKa radiation was used for all XRD studies. A step scan (time constant of one second) was used and the X-ray tube operated at 35 kV and 25 mA. Aluminium sample holders were used and the powdered samples were held in place with a glass slide. Reduced samples were mixed with collodion to prevent oxidation. Silicon powder was used as a standard for crystallite size determination.

4.3.4. Thermal Gravimetric Analysis

Thermal Gravimetric Analysis (TGA) has found widespread application in the study of physical and chemical transitions that involve a weight loss. A notable application is in the study of polymer degradation [227]. 84.

In catalysis, TGA has been used by Shimokawabe et al. [225] in studies of the reduction of copper on silica catalysts prepared by ion exchange techniques, and by Banerjee et al. [153,183] in studies of the preparation of copper chromite catalysts from cupric oxide and chromia. Both studies have made use of the weight loss due to reduction of cupric oxide by the reaction

(2.69)

Complete reduction of CuO results in a weight loss equivalent to 20 percent of the CuO content of the catalyst. Even at low copper loadings (< 1 percent) this weight loss is substantial enough for evaluation by TGA [225].

In this study, a·ou Pont 951 TGA is used to study the rate and degree of reduction of copper chromite catalysts. The instrument consists of a 500 Wresistance-wound furnace and quartz tube enclosing the balance assembly. This assembly is made up of a removable platinum sample pan supported on a counterbalanced quartz rod. Displacement of the rod due to sample weight change is measured with a light beam and photodiodes. Platinel II thermo­ couples are used for temperature control and measurement. Analysis is carried out by zeroing the instrument and loading 10-20 mg of sample onto the pan, the exact weight being obtained when the instrument is re-zeroed. The instrument is then spanned such that the electrical response corresponds to 0-100 percent weight loss on the Y axis of an X-Y recorder. A flowrate of nitrogen or hydrogen (50 cm3min-1) is set and after the quartz tube has been 85.

purged, a rate of temperature increase of 5°C min- 1 is set as the X-axis of the recorder. Weight loss between 50 and 650°C was generally measured although isothermal experiments at 220°C were also carried out. In the latter case, a time base of 0.2 cm min-l substituted for the X-axis. The differential of weight loss, used to accurately determine where a particular period of weight loss had started and finished, was also plotted. Catalyst samples were thoroughly dried at 110°C for 16 hours in an air oven prior to analysis by TGA.

Isothermal TGA at 220°C was also carried out using the apparatus of Figure 4.3. Weight loss was measured by weighing the sample holder before reduction in hydrogen and after reaction with nitrous oxide. 86.

CHAPTER 5

INVESTIGATION OF COPPER CHROMITE CATALYSTS

5.1. Introduction

In Section 2.4.2. it was shown that, in the hydrogenation of vegetable oils, copper chromite catalysts are believed to be reduced by the following sequence of reactions

(2.69)

(2.70)

(2.68)

Reaction 2.69 was believed to occur rapidly accompanied by reaction

2.70. Under high hydrogen partial pressures (> 100 atm), reduction to copper on chromia by reaction 2.68 proceeds to completion, b~t at pressures lower than this, the catalyst is likely to consist of a mixture of copper, cuprous chromite (cu2cr2o4) and chromia.

Reduction conditions employed in this thesis for the pre­ treatment of copper chromite catalysts (Section 4.2.5) differ markedly from those in hydrogenation processes in several important areas:-

1. For the dehydrogenation experiments, catalysts are pre-reduced in a flow of hydrogen. However, during hydrogenation, catalyst reduction accompanies the hydrogenation reaction. 87.

2. Catalysts are reduced in pure hydrogen for the dehydrogenation, and a large part of the catalyst surface would be expected to be covered with adsorbed hydrogen. In hydrogenation, reduction occurs in a liquid medium of oil and solvent (usually methanol or ethanol} and hydrogen adsorption would have to compete with adsorption of these organic species.

3. Reduction is carried out at atmospheric pressures for dehydrogenation, but hydrogen pressures of up to 250 atmospheres are used during hydrogenations.

4. Water liberated by reactions 2.69 and 2.68 is continuously removed in a flow of hydrogen prior to dehydrogenation. However, as hydrogenation is generally carried out batchwise, removal of water from the catalyst surface is not as easily achieved.

For these reasons, there is considerable uncertainty as to the mechanism and extent of reduction during catalyst pre­ treatment for dehydrogenation, and as to the nature of the catalytic surface that results. To resolve this uncertainty elemental analysis, adsorption of nitrogen, reaction with nitrous oxide, X-ray diffraction a~d Thennal Gravimetric Analysis, are used in this section of the research to fully characterise the catalysts before and after reduction. 88.

5.2. Catalyst Characterisation

5.2.1. Elemental analysis

Total copper and chromium content of the unreduced catalysts was obtained using the digestion procedure outlined in Section 4.3.1. The proportion of copper that was present as cupric oxide was estimated by Atomic Absorption Spectroscopy and the extraction procedure of Section 4.2.1. The results are presented in Table 5.1.

When expressed in molar terms, the ratios Cu/Cr of 1.08 and 1.16 obtained for catalyst 1 (1808) and the laboratory prepared catalyst 4 respectively, correspond closely to the value of unity expected for the decomposition of copper anmonium chromate by reaction 2.59.

(2.59)

Analysis of reaction 2.59 suggests that 50 percent of the total copper should exist as cupric oxide, and as shown in Table 5.1, this value is closely approximated for catalysts 1 and 4. There­ fore, it is now assumed that catalyst 1 and the laboratory-prepared catalyst 4 are Type A copper chromites as described in Section 2.4.1. Leaching of the laboratory-prepared catalyst 4 by the procedure of Section 4.2.1. yields a powder with a Cu/Cr ratio of 0.49. This corresponds closely to the stoichiometric value of 0.5 for cupric chromite (Cucr2o4) and indicates that the extraction procedure is virtually quantitative. This powder is now assumed to be pure cupric chromite {Cucr2o4). TABLE 5.1. Catalyst characterization

Composition wt% Surface Area (m2g-1) Crystallite Catalyst Size Chromium Mol. Ratio Total Copper 0 Coeeer Ratio (A) Cu SBET Seu Total As CuO Cr 5cu SBET

1. Harshaw 35.6 18.9 27.1 1.08 28.4 6.8 0.24 80 (1808)

2. Harshaw 59.4 54.5 13.0 3.74 16.3 10.8 0.66 100 (0203)

3. Girdler 33.7 24.5 25.1 1.10 45.1 15.4 0.34 70 (G-22) 4. Laboratory 44.9 24.5 31.6 1.16 40.7 7.9 0.19 110 preparation

5. Cucr2o4 27.3 - 45.1 0.49 59.9 0.9 0.01 30

00 .\0 90.

A Cu/Cr ratio of 1.1 indicates that catalyst 3 {G-22) is also a Type A copper chromite with barium incorporated into the catalyst by the method of Connor et al. [148,156]. However, as some of the chromium is present as barium chromate (BaCr04), an excess of copper {73 percent), as compared with catalysts 1 and 4, is present as cupric oxide.

The preparation of a Type A copper chromite catalyst yields an equimolar mixture of cupric oxide and cupric chromite (Equation 2.59) and therefore an atomic ratio of Cu/Cr approximating unity. Catalyst 2 {0203) is likely to be a Type B copper chromite prepared by the high temperature ignition of cupric oxide and chromia according to the reaction

{2.60) since the Cu/Cr ratio.is far in excess of unity. A large excess of cupric oxide has been used in the preparation which yields a cata­ lyst containing almost 60 weight percent copper. A comparatively small amount of chromia is available for the fonnation of cupric chromite and over 90 percent of the copper remains as cupric oxide.

5.2.2. Measurement of total surface area and pore size distribution

Total surface areas of reduced catalysts, measured by the nitrogen adsorption technique of Section 4.3.2.1., are listed in Table 5.1. A particularly low surface area is obtained for the Type B copper chromite {catalyst 2), yet the surface areas of the commercial Type A copper chromites {catalysts 1 and 3) are still not especially high. The difference in surface area between 91.

catalysts 1 and 3 may be due to the existence of amorphous barium oxide in the latter. Highest surface area is obtained for the

reduced form of cupric chromite {Cucr2o4).

Pore size distributions were determined for catalysts 1 and 2 and the results are shown in Figure 5.1. Distributions are fairly broad, and in the case of catalyst 1, a significant propor­ tion of the surface area {20 percent) is contained in pores with

0 radii in excess of 200 A. However a median in the distributions of

0 50-70 A shows a large contribution to the surface areas of very

0 small pores{< 50 A).

Surface areas before and after reduction are shown for the conmercial copper chromite catalysts in Table 5.2, and a distinct increase in surface area upon reduction is noted.

5.2.3. Copper surface areas

Copper surface areas, obtained by the nitrous oxide tech­ nique of Section 4.3.2.2. , along with the ratio of copper surface area to total surface area, are shown in Table 5.1. Despite low total surface areas and high loadings of copper, copper surface areas are comparatively high and in the case of the Type B copper chromite {catalyst 2), the copper surface area accounts for over 50 percent of the total surface area.

Copper surface areas are similar for two of the Type A copper chromites {catalysts 1 and 4), but the value for the Bao containing catalyst 3 {also Type A) is markedly higher. A very low copper surface area {only 1.5 percent of the total} is obtained on reduction of Cucr2o4• 92.

100 . a: 80 (/")

_J a: I- 60 ~ I-

IJ... ~ 40 :;--.: 20

0 0 100 200 300 400 P~RE RADIUS (ANGSTR~MSl

FIGURE 5.1. Pore size distributions of copper chromite catalysts

TABLE 5.2. Changes in surface area and crystallite size on reduction

Surface Area Crystallite (m2 /g) Size Catalyst -~ngstroms) Before After Reduction Reduction CuO Cu

1808 23.2 29.7 70 80 G-22 45.1 55.2 80 70 0203 8.7 16.3 80 100 93.

TABLE 5.3. Effect of N20 sample size on copper surface area

Copper Surface Area (m 2/g) Catalyst 3.1 mls N2o 1.63 mls N20

1808 7.0 6.8 G-22 16.3 15.4 0203 9.9 10.8

A question arises here as to what extent the high metal sur­ face areas reported in Table 5.1 could be attributed to bulk oxidation of copper by nitrous oxide. Bulk oxidation has been reported to occur above 70°C [211,212], though the results of Evans et al. [216) at 90°C suggest that the overestimation of copper sur­ face area due to bulk oxidation would be less than 10 percent. If no bulk oxidation occurs, then provided an excess of nitrous oxide is used, the copper surface area should be independent of the concentration of nitrous oxide, and duration of the pulse over the sample. If bulk oxidatioo is occurring, then nitrous oxide should be taken up for the duration of the pulse of gas over the catalyst, 94. thus a larger copper surface area should be obtained using a larger sample of nitrous oxide.

In Table 5.3, copper surface areas have been determined for two different volumes of nitrous oxide. A 100 percent increase in gas volume should increase both the concentration of nitrous oxide in the carrier gas, and the duration of the pulse over the catalyst. As shown in Table 5.3, no increase in copper surface area is noted and bulk oxidation is apparently insignificant.

If a second pulse of nitrous oxide was passed over the catalyst 10 minutes after the initial surface area determination, a volume of nitrogen approximately 10 percent of that observed initially, was liberated. This indicates that slow bulk oxidation does occur over time. Bardeen et al. [228] have shown the rate of bulk oxidation to be controlled by the diffusion of cupric oxide, and as this process appears to be very slow compared with the rate of surface reaction of nitrous oxide, the extent of bulk oxidation during the initial pulse is unlikely to be significant.

5.2.4. X-ray Diffraction

Table 5.4 lists literature values of diffraction maxima for copper, cupric oxide, cuprous oxide, cupric chromite, cuprous chromite and chromia. The values for cupric and cuprous chromite have been obtained from the data of Schulz et al. [222] and, as has been mentioned in Section 4.3.3, some uncertainty is involved in interpreting the figures provided. The preparation of a copper chromite catalyst in Section 4.2.1. was carried out with the particular aim of being able to accurately resolve cupric chromite and cupric oxide maxima in the range 35° < 20 < 39°. 95.

TABLE 5.4. Tabulated values for X-ray diffraction maxima 2B

Cu CuO Cu 20 Cucr2o4 Culr2o4 Cr2o3

18.1 24.6 29.5 31.2 31.3 34.6 34.1 35.6 * 36.3 * 35.1 * * 36.1 36.8 * 38.8 37.1 43.2 * 42.2 42.0 40.5 42.3 50 4 49.7 51.6 56.1 55.6 56.2 61.3 61.5 62.0 64.4 64.9 65.0 66.5 74.0 74.3

* Major Peak. 96.

2e 2e

Laboratory-prepared copper chromite

Figure 5.2. XRD spectra.

Figure 5.2. shows part of the XRD spectrum for catalyst 4, the laboratory-prepared copper chromite, and catalyst 5, cupric chromite. The distinct maxima at 35.5° and 38.8° in the spectrum of catalyst 4 correspond closely to the major peaks for cupric oxide in Table 5.2. The XRD spectrum for cupric chromite shows that the peak for cupric chromite as well as that for cupric oxide, and further, that the broad peak at 37.7° is the secondary peak for cupric chromite. Comparing the spectrum for cupric chromite with the data in Table 5.4 shows that there is an error of up to 0.7° in 97. interpretating the figures of Schulz et al. [222]. Accordingly, the spectrum for the leached sample was subsequently used for identification of cupric chromite, the criteria for identification being the existence of the peak at 35.4° if no cupric oxide was present, or the existence of the peak at 37.8° in the presence of cupric oxide. The maxima above 50° and below 35° were too indis­ tinct to be of any value in identification.

The data of Table 5.4 were used for the identification of cuprous chromite (Cu2cr2o4) as it was not known how a pure sample of this compound could be prepared. The major peak at 36.1° could not of course be resolved from the major peak for cuprous oxide, but since the cupric chromite peak at 31.2° is virtually indistinct, the criteria for the identification of cuprous chromite could be satisfied by appearance of the secondary peak in this vicinity as well as the major peak. Further, the existence of cuprous oxide could be discounted if the secondary peak at 42.2° was not identified.

Using these criteria, the composition of the commercial catalysts could be qualitatively assessed. Catalyst reductions were carried out using the apparatus of Figure 4.1 and the pre-treatment conditions outlined in Section 4.2.5.

Figure 5.3 shows XRD spectra for the three commercial copper chromite catalysts before reduction. The major CuO peak (35.6°) is clearly visible, especially for catalyst 2 which has the highest CuO loading. The secondary CuO peak (38.8°) is also very distinct. The secondary CuCr2o4 peak (37.8°) is fairly well defined for 98.

JU ·'le 6E

BE ...... M l,E -

!I:

SE

Cl) +> Cl) ,-~ ,0 +> IO u QJ .,..+> Dh ss.. ..c u r,c r.Ml"llt s.. QJ a. I[ a. 0u ...... 0 N "'C QJ N u - ::, 'I[ "'C QJ 1,'l!,"5[ s.. C: ::, SE ~ 0 IOs.. +> u QJ a. Cl) C 0::: X Uh M. 6( U')

QJ s.. ::, BE .,..C, u.

l,E

!IE

L~"SE S(

h[ 99. catalysts 1 and 2, but no so well defined for catalyst 3. This is most likely due to the incorporation of chromium as barium chromate

(BaCr04) as well as Cucr2o4. The distinct peaks of BaCr04 at 28.1 and 25.3° were evident on the full spectrum of catalyst 3. Figure 5.4 shows the spectra obtained for catalyst 1 and 2 after these catalysts were subjected to the leaching procedure of Section 4.2.1. Quantitative removal of CuO is indicated by the disappearance of the secondary peak (38.8°) for that compound. The secondary peak for Cucr2o4 (37.7°) is now quite distinct, and the contribution made by the primary Cucr2o4 peak (35.3°) to the major peak for CuO (35.6°) is evident. A higher degree of crystallinity for Cucr2o4 in catalyst 2 is suggested by the more sharply defined spectrum.

2e 28

(1) (2)

Figure 5.4 •. XRD spectra of leached copper chromite catalysts. 100.

When the copper chromite catalysts were reduced, the XRD spectra showed the distinct peaks for copper at 43.2° and 50.4°. Crystallite sizes were calculated by the method of Section 4.3.3. 0 and, as shown by Table 5.1, values between 70 and 110 A were obtained. A small peak for copper was also obtained on reduction 0 of Cucr2o4 and the crystallite size of 30 A indicates a high degree of dispersion. Figure 5.5 shows the spectra for reduced catalysts 1, 2 and 3. For the Type B copper chromite {catalyst 2) complete disappearance of CuCr2o4 is evident and a very distinct cu2cr2o4 peak has emerged indicating that transition between Cucr2o4 and cu2cr2o4 has taken place. The very small peak at 39.2° suggests that a small fraction of CuO has remained unreduced. A dramatic decrease in the intensity of the composite Cu0/Cucr2o4 peak is evi­ dent for the Type A copper chromites {catalysts 1 and 3) suggesting that significant reduction of these compounds has taken place. As neither of the secondary peaks for CuO and Cucr2o4 can be distin­ guished, it is impossible at this stage to identify the peak that remains. Peaks for Cu2cr2o4 are barely visible in the spectra for catalysts 1 and 3.

No peaks due to cuprous oxide {Cu20), cupric chromate {CuCr0 ) or chromia were evident in the XRD spectra. Chromia formed 4 . during catalyst reduction is likely to be amorphous [162]. The peaks evi~ent for BaCr04 in the spectrum of unreduced catalyst 3 disappear completely on reduction, presumably due to the formation of amorphous barium oxide {BaO) by the reaction

{5.1) ,- I ~ iii I!! ,.,,"-: l"!"I

ii! ~-- J \ p "' iii

2e 2a 2e

~ II: Ill i.: IR II: ~ II: Ill ~ IR II: §I '1 T ,.,, UI C'" m m m m

(1} . (2} (3}

Figure 5.5. XRDspectra of reduced copper chromite catalysts.

..... 0 ..... 102.

In Table 5.2, crystallite size of cupric oxide (CuO) in the unreduced corrmercial catalysts is compared with the corresponding copper cystallite size obtained from the spectra of the reduced catalysts. In each catalyst, crystallite sizes are similar for the two fonns of copper.

5.2.5. Thermal Gravimetric Analysis

Initially, TGA was carried out on the copper chromite cata­ lysts over the temperature range 50 to 650°C under a hydrogen atmosphere. The weight loss profiles that were obtained are shown in Figure 5.6. Two distinct phases of weight loss, a sharp loss between 200 and 250°C, followed by a more gradual loss up to 550°C, are apparent. Only the second phase of weight loss is evident for cupric chromite.

The weight loss due to reduction of CuO can be predicted from the stoichiometry of Equation 2.69.

{2.69)

As shown in Table 5.5, the predicted values and the observed values for the first phase of weight loss in Figure 5.6, are very similar. It would appear then, that the first stage of weight loss is due to reduction of CuO by reaction 2.69. Thus the broad peaks at c a. 35.5° apparent in the XRD spectra (Figure 5.5) of reduced catalysts 1 and 3, are due to unaltered Cucr2o4• Attempts were then made to relate the weight loss up to 650°C to reduction of the total copper in Table 5.1, and as shown in Table 5.5, good correlation between predicted and observed values is evident only 103.

0 L Catalyst 1 8 '-- 10 12 ----- 0 L 8 \_ 10 Catalyst 2 12 .,, 0 ' .,, L 0 8 Catalyst 3 ..J "-- 10 I- J: 12 Cl ---- 0 w L ~ '\._ 8 Catalyst 4 ~ . 10 12 ---- 0 L Cucr2o4 8 10 ------12

0 100 200 300 LOO SOO 600

T E M p E R A T u R E I "c I

FIGURE 5.6. Weight loss profiles for reduction of copper chromite catalysts. From top: Catalysts 1, 2, 3, 4 1 Cucr2o4 104.

TABLE 5. 5. Observed and predicted weight losses due to catalyst reduction (wt%) Temperature 50-650°C 220°c

Reduction of CuO Reduction of Catalyst Total Cu Obs Pred 06s Prea Obs

1 4.5 4.76 10.8 8.96 4.8 2 13.2 13.71 15.2 14.97 12.9 3 6.2 6.17 12.6 8.50 7.5 4 6.2 6.17 11.6 11. 31 6.0

5 9.0 6.86 < 1.0

for catalysts 2 and 4. The discrepancy in the case of catalysts 1 and 3 is probably due to decomposition of binder and other residues incorporated during the catalyst preparation. When TGA of the copper chromite catalysts was carried out under nitrogen, weight losses equivalent to the discrepancies in Table 5.5 were obtained (Table 5.6.)

TABLE 5.6. Weight loss by TGA for copper chromite catalysts­ nitrogen atmosphere, 50-650°C Weight Loss Discrepancy in Table 5.4 Catalyst (wt%) (wt%) 1 (1808) 2.0 1.8 2 (0203) 0.2 0.2 3 (G-22) 3.9 4.1 4 (Lab. prep.) 0.5 0.3 105.

The magnitude of these weight losses is reflected by the concen­ tration of non-catalytic components (binder etc.) in the catalysts (Table 4.4).

To obtain a more reliable measurement of the weight losses expected during catalyst pre-treatment, isothermal TGA, under hydrogen at 220°C, was carried out for each catalyst over 16 hours. Only the first stage of weight loss was evident, and only minimal weight loss was recorded for Cucr2o4 (Table 5.5). Except for catalyst 3, the magnitude of the weight loss at 220°C corresponds closely to the values predicted by Equation 2.69, indicating that under the conditions employed during catalyst pre-treatment, only reduction of CuO would be expected. In the case of catalyst 3, the extra weight loss is probably due to formation of BaO by reac­ tion 5.1 which predicts a further weight loss of 1.7 percent given the nominal 10 percent BaO content specified {Table 4.4). The TGA profile obtained for the isothermal reduction of catalyst 3 showed the initial phase of weight loss due to reduction of CuO, to be followed by a gradual weight loss of approximately 2 percent proceeding over 100 minutes. This is in agreement with the find­ ings of Russian authors [180] who have shown decomposition of

BaCr04 to occur at a much slower rate than reduction of CuO.

A disadvantage of the TGA technique described thus far, is that hydrogen is not directed through the catalyst sample but rather, takes a path of less resistance around the sample pan, and so the mass transfer conditions of a tubular reactor, as employed for activity testing, are not fully satisfied. Water evolved by reaction 2.69, would be continuously removed by the flow of 106. hydrogen through a fixed bed of catalyst in the tubular reactor, but may stay adsorbed on the catalyst surface during TGA thus inhibiting further reduction. Accordingly, isothennal TGA, using the apparatus of Figure 4.3 which incorporates a fixed bed reactor, was carried out on the copper chromite catalysts to be used in activity measurements in later chapters. Weight losses were measured gravimetrically and are compared in Table 5.7 with the observed and predicted values of weight loss reported in Table 5.5. No further weight loss is evident for the Type B copper chromite (catalyst 2), but the use of a tubular reactor does appear to facilitate further weight loss in the case of the Type A copper chromites (catalysts 1 and 3).

TABLE 5.7. Predicted and observed weight losses for catalyst reduction at 220°c

Catalyst Weight Loss (%} Predicted For DuPont TGA For Tubular Reactor

1. 1808 4.8 4.8 5.7 2. 0203 13.7 12.9 13.0 3. G-22 6.2 7.5 9.0

5.3. Discussion

Characterisation by elemental analysis and nitrogen adsorp­ tion (Table 5.1), shows the unreduced copper chromite catalysts to contain high loadings of copper(> 30 weight percent) and to be of comparatively low total surface area. This combination of physical characteristics would be expected to yield, on reduction, a poorly 107. dispersed copper catalyst with low metal surface area, especially as copper is considered to be an easily sintered metal. Instead, a relatively high degree of dispersion, as indicated by high metal surface areas and low copper crystallite size, is apparent from the results in Table 5.1. Aggregation of copper on reduction is not apparent and total surface area actually increases slightly (Table 5.2). Further, although pore size distributions (Figure 5.1) indicate a high proportion of the total surface area to be 0 contained in very small pores(< 50 A), the copper crystallite sizes of Table 5.1 could only be contained in comparatively large pores. This result is significant as it shows that relatively high metal dispersion can be obtained without distribution of the metal in very small pores where mass transfer considerations may become significant.

Characterisation by XRD and TGA gives added insight into the nature of the reduced catalyst and the mechanism of reduction. The XRD spectrum of the Type B copper chromite (catalyst 2) shows distinct maxima for cuprous chromite (cu2cr2o4) which can form by reduction of cupric chromite (Cucr2o4) by the reaction [172]

(5.2) or by a ·reverse disproportionation reaction between copper metal and Cucr2o4

(2.70) 100.

Reaction 5.2 requires a weight loss due to reduction of CuCr2o4 as well as the weight loss that is attributed to reduction of CuO by reaction 2.69. However, TGA of catalyst 2 shows that no further weight loss is apparent after reaction 2.69 has proceeded to completion (Tables 5.5, 5.7). In any case, Stroupe [172] has shown that very high temperatures(> 900°C) are required to effect reaction between Cucr2o4 molecules presumably as the spinel struc­ ture of cupric chromite must be broken down.

Formation of cu2cr2o4 appears much more likely to take place via reaction 2.70 and to involve diffusion of copper atoms, formed by reaction 2.69, to the cupric chromite spinels. As the XRD spectra for cupric and cuprous chromite are similar (Table 5.4), it would seem that fonnation of cu2cr2o4 does not involve break­ down of the spinel structure, however it is very difficult to speculate on the exact nature of this reaction. The formation of cu2cr2o4 is consistent with XPS studies of copper chromite cata­ lysts which show that a high surface concentration of monovalent copper is formed after the first few minutes of reduction [162,178].

Cuprous chromite should also be fanned on reduction of the Type A copper chrornites (catalysts 1 and 3) as copper and Cucr2o4 would be certain to be in as intimate contact in these catalysts as in catalyst 2. However, the XRD spectra obtained when the catalysts are reduced, barely show the presence of cu2cr2o4, though a significant decrease in the intensity of the CuCr2o4 peaks is noted. This apparent inconsistency can be explained by the results of Table 5.8 which show a weight loss on top of that expected by the reduction of CuO, when TGA is carried out with a 109.

tubular reactor. This indicates that cu 2cr2o4 has been further reduced to copper and chromia by reaction 2.68

{2.68)

The increase in surface area observed in Table 5.2 could be due to the formation of amorphous chromia by this reaction. Although some minor reduction of pure Cucr2o4 is evident in Table 5.1, XPS [162,178] has shown that after reduction of CuO, no Cu{II) is available on the surface for further reduction, and so the contri­ bution to weight loss of reduction of Cucr2o4 by the reaction

{5.3) is likely to be minimal.

So, while the Type B copper chromite {catalyst 2) reduces to copper supported on cu2cr2o4; the Type A copper chromites {catalysts 1 and 2) appear to reduce to copper on chromia, with amorphous chromia presumably isolating a core of Cucr2o4 that has not taken part in reaction 2.70, from further reaction. No unreacted Cucr2o4 is apparent on reduction of catalyst 2 as there is a considerable excess of copper oxide to form cu2cr2o4 by reactions 2.69 and 2.70. The large excess of copper may also explain why catalyst 2 does not undergo further reduction by reaction 2.68 even when a tubular reactor is employed {Table 5.7), as copper could effectively shield cu2cr2o4 from further reaction. A very high ratio of copper surface area to total surface area {0.66) is 110. evident for this catalyst in Table 5.1. The higher degree of crystallinity apparent for CuCr2o4 in catalyst 2 (Figure 5.4) may also serve to retard further reduction.

An important function of reaction 2.70 may be to disperse the copper formed by reduction of CuO, and in doing so prevent the aggregation of large copper crystallites on the catalyst surface. Thus the relatively high degree of metal dispersion evident from the data of Table 5.1, may be strongly related to the mechanism of catalyst reduction.

Finally, the elemental analysis of Section 5.2.1 can be combined with the TGA data of Table 5.8, to allow calculation of weight fractions of copper, Cucr2o4, Cu2cr2o4 and Cr2o3 in each conmercial catalyst after reduction by the pretreatment of Section 4.2.5. Calculations are given in Appendix V and the results presented in Table 5.8.

TABLE 5.8. Composition of copper chromite catalysts after reduction

Composition (wt %)1. Catalyst Cu Cucr2o4 Cu2tr2o4 Cr2o3

1808 23.9 50.7 8.6 27.2 G-22 2• 31.6 20.0 0203 50.2 42.3

1. Balance= Inerts 2. 12.2 percent BaO 111.

5.4. Conclusions

Despite having the characteristics of low total surface area, and high copper loading, the three corrmercial copper chromite catalysts to be used for methanol dehydrogenation have been shown to reduce to relatively highly dispersed copper catalysts as indi­ cated by high metal surface areas and small copper crystallite size.

Reduction is believed to occur in three stages designated by the reactions

(2.69)

(2.70)

(2.68)

Reaction 2.69 occurs rapidly in pure hydrogen at 220°C, accompanied by formation of cuprous chromite by reaction 2.70. Type A copper chromites were further reduced by reaction 2.68 to yield a cataly­ tic surface of copper supported on chromia, though some unreduced cupric chromite remain,.. The large excess of copper, present on reduction of the Type B copper chromite, effectively shields cuprous chromite from further reduction by reaction 2.68 and a catalytic surface of copper supported on cuprous chromite results.

These findings confirm the results of earlier studies where XPS was used to follow catalyst reduction during liquid phase hydrogenation. 112.

CHAPTER 6

PRELIMINARY INVESTIGATIONS

6.1. Introduction

As experience within this laboratory had shown copper chromite catalysts to be highly effective for the hydrogenolysis of methyl formate (Section 2.3.4.), these catalysts were the obvious choice for studying the feasibility of methanol dehydrogenation to methyl formate. Accordingly the three co111nercial copper chromite catalysts studied in Chapter 5 were used to establish a set of reaction conditions where comparative studies with other catalysts could be carried out. It was also hoped to be able to explain differences in activity and selectivity within the copper chromite catalysts on the basis of the results of catalyst characterization in Chapter 5.

6.2. Thermodynamic Yields of Methyl Formate

To establish the conditions of pressure, temperature and concentration that might be used to study methanol dehydrogenation over copper chromite catalysts, theoretical equilibrium yields under different reaction conditions were calculated from thermo­ dynamic data. Details of the calculations are given in Appendix VI along with the BASIC computer program used to solve the thermo­ dynamic equations. Results are presented here in graphical form.

Figure 6.1 shows the effect of temperature on the thermo­ dynamic yield of methyl formate at 1 atm. Obviously, very high 113.

a ,.,,a

a co N

a (0 N

a ,q N

a N N

a ,..... a ...... ,u N w a Q'.'.'. co ::J t- - CI Q'.'.'. a w (0 a.. '.L - w t- a -,q a N- a a- a a a a a a ,q ,.,, (0 l/) N - a [ /. ) N~ISH:3/\NrJJ

FIGURE 6.1. Effect of temperature on thennodynamic yield. P = 1 atm, YCH OH= 1 3 114.

temperatures would be needed to achieve high conversions (> 90 per­ cent). However, as shown in Table 2.1, formation of methyl formate from formaldehyde starts to become thermodynamically unfavourable above 300°C. A temperature range of 180-300°C, as suggested by Patent literature [6], was chosen for investigation.

Figure. 6.2 shows the effect of pressure on thermodynamic yield at 220°C. As would be expected from the stoichiometry of Equation 1.2, a decrease in total pressure increases yield. Clearly, nothing is to be gained from operation of the reaction at pressures in excess of one atmosphere, and although kinetic studies have been performed at lower pressures [122], a reaction pressure of one atmosphere is most practical from an industrial standpoint. Accordingly kinetic measurements were conducted at a nominal pressure of 1 atmosphere. A pressure drop across the reactor of up to 0.3 atmospheres was encountered at the highest space velocity employed. However, as can be seen from Figure 6.2, the effect on equilibrium yield is negligible.

The effect of methanol concentration on thermodynamic yield at 1 atmosphere and 220°C is shown in Figure 6.3. A decrease in methanol concentration increases yield but a 5:1 dilution of methanol is required for the increase to become significant. For this reason, and because one of the advantages of the process outlined in Chapter 1 was the use of a single component feed, dilution of the methanol feed was not normally carried out. In any case, dilution with nitrogen or helium would necessitate chromatogr~phicseparationbetween hydrogen and carbon monoxide for selectivity to be determined (Section 4.1.4). This would not have 115.

70

60 ~

50 -~ • -~ -0 z 40 ~- D Vl ~. a.: UJ > 30 z D u ...... 20

10

00.1 1 10 PRESSURE ( Atm.)

FIGURE 6.2. Effect of pressure on thennodynamic yield 220°C, YcH OH= 1 3 116.

60

50

~ ~ .,__, 40 z D t--1 30 en 0:: w > z 20 D u 10

0 0 . 2 . 4 . 6 . 8 1 1CCH3LJHl

FIGURE 6.3. Effect of methanol concentration on thermodynamic yield. 220°c, P = 1 atm

1 1

I-' I-'

I-' I-'

--... --...

8 8

. .

.6 .6

.4 .4

'J'(HC~~CH3) 'J'(HC~~CH3)

.2 .2

methyl methyl

and and

0 0

10 10

0 0

20 20

30 30

40 40

50 50

thermodynamic thermodynamic

on on

D D u u

z z

> >

w w

en en

n::: n:::

0 0

z z

1--4 1--4

,__, ,__,

X X

,...... ,......

atm atm

1 1

concentration concentration

1 1

= =

P P

hydrogen hydrogen

8 8

. .

220°c, 220°c,

concentration concentration

of of

formate formate

Effect Effect

yield. yield.

6.4. 6.4.

.4 .4 .6

FIGURE FIGURE

'J'(H'J'OR~GEN) 'J'(H'J'OR~GEN)

.2 .2

0 0

10 10

o o

20 20

40 40

30 30

50 50

u u

0 0

z z

> >

w w

n::: n:::

(f) (f)

D D

z z

1--4 1--4

,__, ,__,

X X ,...... ,...... 118. been achieved easily and would have considerably complicated the chromatographic analysis of Section 4.1.2. Dilution with hydrogen is out of the question due to the strong inhibiting effect of hydrogen on thermodynamic yield as shown in Figure 6.4. A somewhat lesser inhibiting effect is observed for methyl fonnate (Figure 6.4).

Thus, apart from flowrate, the only parameter that can realistically be varied to affect equilibrium yield is reaction temperature. Experiments were now carried out at atmospheric pressure, 180 to 300°C and by varying the LHSV of the pure methanol feed over the range 12 h-l to 200 h- 1•

6.3. Results

6.3.1. Activity and selectivity of copper chromite catalysts

In Figures 6.5, 6.6 and 6.7, conversions and selectivities are plotted against LHSV for the experimental results at 200, 220 and 240°C. At 180°C, conversion with the most active catalyst was only 5 percent at an LHSV of 10 h- 1• Above 240°C, conversion continued to increase but the decrease in selectivity was such that the outlet concentration of methyl formate increased only minimally, and began to fall at 280°C.

The codes for catalyst identification in Figures 6.5, 6.6 and 6.7 are:- 0 Harshaw 0203 D Girdler G-22 l:l Harshaw 1808 119.

so -;--! ...... , z 40 ID -(/) 0::: w 30 > z ID u 20

10

0 0 50 100 150 200 100

-;--! - 90 >- ...... -> 80 ...... -u w _J 70 w (/) 60

50 0 50 100 150 200

LHSV ( I /HR)

FIGURE 6.5. Effect of LHSV on conversion and selectivity at 200°( 120.

50 r--, :,...: ....., z 40 D -(/) 0:::: w 30 > z D u 20

10

0 0 50 100 150 200 100

-:--: ...... 90 >- t- -> BO t- -u w .....J 70 w (/) . 60

50 0 so 100 150 200 LHSV ( /HRl

FIGURE 6.6. Effect of LHSV on conversion and selectivity at 220°c 121.

50 X z 40 0 ...... en a::: w 30 > z 0 u 20

10

0 0 50 100 150 200 100

,..... X ....., 90 >- (!] r ...... > 80 ...... r u w _J 70 w (/) 60

50 0 50 100 150 200 LHSV ( 1 /HR)

FIGURE 6.7. Effect of LHSV on conversion and selectivity at 240°( 122.

As would be expected, Figures 6.5, 6.6 and 6.7 show an inverse relationship between LHSV and conversion and, as shown in Figure 4.2, the conditions of a differential reactor are approxi­ mated at low conversions (< 7 percent). Catalysts 1808 and G-22 show similar activity while catalyst 0203 is markedly more active, but the difference is only significant above 200°C and at conver­ sion levels in excess of 10 percent.

Yields in excess of 75 percent of the thermodynamic values of Figure 6.1 can readily be obtained at low ~HSV's, but above 200°C, selectivities fall off dramatically when LHSV falls below 50 h- 1. The main cause of the lower selectivities is a sharp increase in levels of carbon monoxide. Carbon dioxide levels never exceed 0.5 percent. Selectivity differences between catalysts follows similar patterns to conversion differences. At 200°C, the differences are not significant, but above this temperature, catalyst 0203 appears markedly more selective than the other two. Apart from minute traces of methane{< 0.01 percent) no other reaction products were apparent at any stage. Water levels {0.1- 0.2 percent) did not vary appreciably with LHSV or temperature thus dehydration activity was insignificant.

6.3.2. Catalyst stability

Stability of each catalyst was tested at 220°C and an LHSV of 26 h-1• As Figure 6.8 shows, steady deactivation {ea. 20 per­ cent}of catalyst 0203 occurs over 200 minutes. Catalysts G-22 and 1808 show stable activity over the same period though some deactivation is evident after 1000 minutes on line. 123.

0 0 N

0 -l[)

U) 0 w 0 ._ :::> - z :E -....., w :E -._ I') N Q) 0 N 0 0 N Q) l[) 0 C) ...

~ El '4

0

0 0 N ..... 0

( t. l NQ I S~3t\NQJ

FIGURE 6.8. Stability of copper chromite catalysts. 220°C, LHSV = 26 h-1 124.

Table 6.1 shows the results of various attempts to restore the activity of catalyst 0203. At 220°C, conversion was reduced from an initial value of 26.4 percent to 22.7 percent after a methanol flowrate had been maintained at an LHSV of 26 h-l for 300 minutes. Activity was only partially restored when the catalyst was flushed with hydrogen over an extended period, even when the temperature was raised to 300°C. Complete reactivation was achieved by flushing with a mixture of 1 percent o2 in helium over 16 hours followed by reduction, although a temperature of 300°C was required.

TABLE 6.1. Activation/deactivation of catalyst 0203

Testing Conditions Conversion ( %)

Fresh catalyst 26.4

After LHSV = 26 for 300 minute5 22. 7 ' 3 After H2 for 2 h (100 cm /min, 220°C) 22.9 3 After H2 for 60 h (100 cm /min, 220°C) 23.7 3 25.1 After H2 for 2 h (100 cm /min, 300°C) 3 After 1% o2 for 16 h (100 cm /min, 300°C) 26.5 3 After 1% o2 for 16 h (100 cm /min, 220°C) 23.7 125.

6.4. Discussion

The data of Figures 6.5, 6.6 and 6.7 were obtained with the aim of establishing conditions for the comparison of copper chro­ mite catalysts with other catalysts for the dehydrogenation of methanol to methyl formate. As can be seen from Figure 6.5, 200°C is not a particularly suitable temperature as the activity and selectivity differences between the catalysts are not great and conversions are fairly low. At 240°C {Figure 6.7), the differences are more marked, but fairly high flowrates {LHSV > 200 h-1) would be needed to achieve conversions low enough for the initial rate approach to be valid. At 220°C {Figure 6.6), high conversions, where differences in conversion and selectivity are obvious, and low conversions suitable for the initial rate approach can readily be obtained within the range of flowrates available. At this temperature, equilibrium conversion of methanol to methyl formate is approximately 40 percent.

Having established 220°C to be the most appropriate tempe­ rature for later catalyst comparisons, a number of important conclusions can be drawn from the results of Section 6.3. Firstly, as no fonnaldehyde was observed, the rate of formaldehyde condensation to methyl formate by the reaction

2HCHO + HCOOCH3 {2.33) is undoubtedly very rapid and the rate controlling step is most likely to be the dehydrogenation of methanol to fonnaldehyde.

{2.28) 126.

Further, if the initial rates. obtained from Figures 6.5, 6.6 and 6.7 by the method outlined in Section 4.1.4 .• are fitted to an Arrhenius relationship, straight lines are obtained and an apparent activation energy of ea. 80 kJ mole-l is obtained for each catalyst (Figure 6.9). This indicates that mass tran5fer control is not significant and the rate of surface reaction of Equation 2.28 was being observed.

12

• 0203

l 1 [!] G 22 w A 1808 t- a: a::: 10 CJ ~ __J I 9

B l.S l • SS 2 2 • 0 5 2 . l 2.15 l /T X 1 OOO ( 1 /K J

FIGURE 6.9. Arrhenius plot for copper chromite catalysts 127.

Comparison of activity and selectivity between the copper chromite catalysts also merits discussion. In Figure 6.10, conversions obtained from Figures 6.5, 6.6 and 6.7 for an LHSV of 25 h-1 are plotted for each temperature, and clearly, catalyst 0203 is markedly more active than the other two catalysts. This result is surprising as catalyst G-22 has the highest copper surface area (Table 5.1) and would be expected therefore to have the highest activity. This is assuming, of course, that chromia has no dehydro­ genation activity and, as will be shown in Chapter 7, this is indeed the case.

50

,...... ~ 0203 ~....., 40 z [!) G 22 0 30 1808 -en • er: w > z 20 0 u 10

0 200 220 240 TEMPERATURE (Cl

FIGURE 6.10. Effect of temperature on conversion LHSV = 25 h-1 128.

The apparent discrepancy in catalyst activity can be explained by pre-empting the results of Chapter 7 which will show a marked effect of chromia in supressing specific activity for dehydrogena­ tion. Clearly, activity will be affected by the amount of chromia in the reduced catalyst. As Table 5.8 shows, catalyst 0203 has no chromia present after pretreatment, and it is not surprising then that it is the most active.catalyst. Further, if the support does interact with copper to reduce specific activity, the effect will be magnified as dispersion of copper is increased. Boudart [229] has proposed that the modification of specific activity will be 0 most pronounced if crystallite size is less than 50 A. Dispersion {D} is calculated from the following relationship [230]:

D =no.of surface metal atoms total no. of atoms {6.1}

The number of surface metal atoms can be calculated from copper surface area {Table 5.1} by assuming a surface density of copper of 1.46xlo19 atoms m-2 [216], and the total number of atoms from the results of Table 5.8. Dispersions for each catalyst then are:-

0{0203} = .032

0(1808} = .043 D{G-22} = .073

A sample calculation is given in Appendix VII.

Clearly, despite similarities in crystallite size and copper loading {Table 5.1}, catalyst G-22 is markedly more dispersed than catalyst 1808. A plausible reason for this effect is that the 129. reduction of cuprous chromite to copper and chromia by the reaction

(2.68) appears to have proceeded to a greater extent in catalyst G-22, as evidenced by higher levels of chromia (Table 5.8). Johansson and Lundin [162] have shown that copper formed by Equation 2.68 is amorphous, thus despite the similarity in observable crystallite sizes, amorphous copper makes a greater contribution to dispersion in catalystG-22than in ~atalyst 1808. The activity of catalyst G-22 is affected then by both the presence of chromia and the dispersion of copper on the chromia support. These two effects counteract the influence of high surface area thereby reducing catalytic activity.

An added consideration may be the presence of barium oxide. It will be shown in Chapter 7 that magnesium oxide also reduces specific activity, and as the catalytic properties of magnesium and barium oxides are very similar (231] due to the similar posi­ tions of the metals in the Periodic Table, it would not be unexpected that the presence of b~rium oxide could also reduce the activity of catalyst G-22.

Selectivity differences are also apparent amongst the three copper chromite catalysts. Selectivities obtained from Figures 6.5, 6.6 and 6.7 for an LHSV of 25 h-l are plotted in Figure 6.11 for each temperature and clearly show catalyst 0203 to be markedly more selective. Decreases in selectivity arise from the formation of carbon monoxide and carbon dioxide, the latter being derived 130.

100

,-. ;,-.: ...... , 90 >- I- -> 80 I-- u w _J 70 ~ 0203 w <.n C!1 G 22 60 • 1BOB

50 200 220 240 TEMPERATURE (Cl

FIGURE 6.11. Effect of temperature on selectivity. LHSV = 25 h-1

from trace water by the steam reforming of methanol outlined in Section 2.3.3. As has been mentioned, carbon dioxide levels never exceeded 0.5 percent, and the formation of carbon monoxide at levels up to 3 percent had the greatest effect on selectivity.

In section 4.1.4, calculation of selectivity was based· on the assumption that the catalysts were active for the decarbonyla­ tion of methyl formate according to the reaction 131.

(2.12) as the formation of carbon monoxide by methanol decomposition according to

(2.31) was not catalysed by copper below 300°C (Section 2.2.3). To test the activity of the catalysts for reaction 2.12, an undiluted flow of methyl formate was passed over catalyst 1808 at an LHSV of 13 h-l while the reaction temperature was increased from 180°C to 300°C. In Figure 6.12, the mole fractions of carbon monoxide and methanol that were obtained are plotted against temperature after having made the appropriate corrections for methanol impurity (c.a. 3 percent) in the methyl formate. As can be seen from Figure 6.12, the copper chromite catalyst was highly active for methyl formate decarbonylation producing equimolar quantities of carbon monoxide and methanol below 240°C, as would be expected from the stoichio­ metry of Equation 2.12. Above 240°C, the yield of carbon monoxide increases much more rapidly than the yield of methanol presumably as the higher temperatures are able to overcome the thermodynamic limitation to methanol dehydrogenation by the excess methyl formate.

Clearly, catalyst selectivity will be dominated by the ability of the catalyst to promote the decarbonylation of methyl formate. Higdon et al. [44] have demonstrated the ability of a wide range of metal oxides including chromia, to catalyse the decarbonylation of methyl formate. The activity of chromia for the reaction will be further demonstrated in Chapter 7. It would 132 .

• 5

z . 4 X CH30H D t- -u a: .3 (!) co 0::: LL w _j • 2 ,D ~

• 1

0 180 200 220 240 260 280 300 TEMPERATURE (Cl

FIGURE 6.12. Decarbonylation activity of copper chromite 1808. YHCOOCH = 1, LHSV = 13 h-1 3

be expected that catalysts G-22 and 1808 which contain chromia should have lower selectivity than catalyst 0203 which contains no chromia at all (Table 5.8). As can be seen from Figure 6.11, this is in fact the case. That catalyst 0203 yields carbon monoxide at all can be related to a distinct activity of copper for decarbony­ lation that will be demonstrated in Chapter 7.

The presence of chromia may also have a profound effect on catalyst stability. As shown in Figure 6.8, steady deactivation 133. of catalyst 0203 is observed over a period of 200 minutes. de Pinillos and Victor [45] have also observed deactivation of copper chromite catalysts for methanol dehydrogenation and Chono and Yamamoto [6] have commented on the deactivation, in reducing atmospheres, of copper catalysts in general. Deactivation has been observed in methanol dehydrogenation to fonnaldehyde by Gosh et al. [98,99,100], and in ethanol dehydrogenation by Dolgov et al. [78,80] and Franckaerts and Froment [82]. In the latter case, deactivation was attributed to polymerization of acetaldehyde on the basis of the findings of Stegner et al. [79] who studied the deactivation of a copper catalyst for ethanol dehydrogenation.

The deactivation of catalyst 0203 could not be reversed by flushing with hydrogen or helium at temperatures as high as 300°C, even over very long periods of time {Table 6.1). This result rules out the possibility of deactivation being due to strong chemisorption of formaldehyde, water or methyl fonnate. As only one catalyst is affected, and an excess of hydrogen is always present, catalyst oxidation is not a possibility. As the mechan­ isms of ethanol and methanol dehydrogenation are similar, it would not be unexpected that the processes giving rise to deactivation would be similar also. The fact that catalyst activity could be restored by flushing at 300°C with 1 percent o2 certainly points to oxidation of high molecular weight carbonaceous products.

Walker [190] has reviewed the mechanism of fonnaldehyde polymerization which can be sunmarized as

(6.2) 134.

A degree of polymerization (n) as low as 100 can yield a material with a melting point of almost 200°C [190]. A wide variety of materials have been suggested to catalyse the polymerization including sulphur, peroxides, organometallic compounds and methyl formate [190]. In the experiments to be carried out with formal­ dehyde vapour in Chapter 9, formaldehyde polymer was observed to form very readily in the reactor tube and ancillary lines.

Since acetaldehyde polymerization has been shown to cause deactivation of copper catalysts for ethanol dehydrogenation, and since formaldehyde polymerizes readily to solids of low vapour pressure, the deactivation of catalyst 0203 will be attributed to the blocking of active sites by the growth of polymeric species derived from the formaldehyde intermediate of the reaction mechan­ ism. Although no research has been carried out into formaldehyde polymerization on copper catalysts, Riekert [232] has shown how polymerization of a slight fraction of formaldehyde causes deacti­ vation of platinum catalysts for formaldehyde decomposition.

Describing the mechanism of catalyst deactivation does not however explain why catalyst 0203 deactivates at a much more rapid rate than the other two copper chromite catalysts. Again the solution to the difference in behaviour of the catalysts appears to be related to the presence of chromia. Church and Joshi [81] have shown how the presence of chromia stabilizes copper catalysts against deactivation during ethanol dehydrogenation to acetalde­ hyde. Moorjani et al. [86] have noted similar results and attributed the effect of chromia to its ability to promote the breakdown of carbonaceous species. It is proposed then that the 135.

presence of chromia in catalysts G-22 and 1808, as well as promo­ ting the decarbonylation of methyl formate, also promotes the decomposition of formaldehyde polymers, thus prolonging the activity of the catalysts. Catalyst 0203 contains no chromia (Table 5.8) and so deactivation occurs much more readily. Furthermore, noting that this catalyst has the lowest total surface area, the effect of polymerization on activity will be more pronounced than if active sites were distributed over a larger catalytic surface.

6.5. Conclusions

The dehydrogenation of methanol to methyl formate occurs readily and with high selectivity over copper chromite catalysts in the temperature range 200°C to 240°C, the most favourable tempe­ rature for catalyst comparison being 220°C. The rate of production of methyl formate appears to be limited by the rate of methanol dehydrogenation to formaldehyde.

Differences in activity, selectivity and stability between copper chromite catalysts can be explained via the results of cata­ lyst characerization in Chapter 5. The presence of chromia, which appears to reduce activity and selectivity but prolongs catalyst lifetime, is a crucial factor in the suitability of the catalysts for methanol dehydrogenation. 136.

CHAPTER 7

PROPERTIES OF COPPER-BASED CATALYSTS FOR METHANOL DEHYDROGENATION

7.1. Introduction

In Chapter 6, test conditions were established for comparing the activity and selectivity of catalysts for the dehydrogenation of methanol to methyl formate. Although copper chromite catalysts were effective for the reaction, a detrimental effect of chromia on activity and selectivity was suspected, and it was hoped that a superior catalyst could be found by screening a wide range of catalysts and supports. Further, by comparing the specific activ­ ity of the catalysts, added insight into the nature of dehydrogena­ tion reaction might be obtained. As the reaction for steam reforming of methanol (Equation 2.50) is believed to proceed via methanol dehydrogenation (Section 2.3.3.), particular attention was paid to those catalysts that were known to be effective for the steam reforming reaction.

This part of the thesis details the preparation of catalysts for the screening experiments, and the results of those experiments. Raneycoppercatalysts were initially prepared as part of this study~ However, the results that were obtained differed markedly from those obtained with the other copper catalysts, and will be discussed separately in Chapter 8. 137.

7.2. Preparation and Characterization of Catalysts

7.2.1. Copper powder

A poorly dispersed catalyst of pure copper, with no possi­ bility of support interaction, was obtained by reduction of cupric oxide powder (Ajax Chemicals 99.0 percent). The powder was diluted with 0.5 g of silica to reduce the pressure drop across the reactor. Metal surface area and crystallite size of the powder are shown in Table 7.2. Dispersion, estimated from metal surface area by the procedure of Appendix VIII, was equal to 0.001. Average particle size, as measured with a Micromeritics Sedigraph, was approximately 10 microns.

7.2.2. Copper supported on silica

7.2.2.1. Impregnated catalysts

Catalysts of copper supported on silica, an inert support, have proven highly effective for the steam reforming of methanol (Equation 2.50) and so were an obvious choice for the screening experiments [139,140,141]. The impregnation procedure of Section 4.2.2. was used to prepare catalysts of varying copper content and metal surface area. Each catalyst was fully characterized and the results shown in Table 7.1.

Metal surface area increased with copper content up to a loading of 15.6 weight percent. Crystallite size increased also indicatingagreater extent of aggregation at higher surface concentrations of copper. Above 15.6 weight percent, the extent of aggregation is such that metal surface area decreases slightly. A similar trend has been found by Hassan et al. [206] for 138.

TABLE 7.1. Characteristics of copper/silica catalysts prepared by impregnation* Surface Area (m2g-1) Copper Crystallite Dispersion Content size (wt%) Total Copper (A)

26.5 183 3.6 240 0.021 15.6 227 4.8 230 0.048 9.1 240 1.6 180 0.027 4.4 261 0.9 190 0.032 2.0 272 0.4 140 0.031 1.0 285 0.4 110 0.062

* Silica support of surface area 322 m2g-1• Code ID 57 of Davison Chemical Division of W.R. Grace and Co. increased loadings of copper supported on alumina. The crystal­ lite sizes that were obtained were surprising as characterization of the silica support by the nitrogen adsorption/desorption technique of Section 4.3.2.1. showed that almost 100 percent of the support surface area could be attributed to pores with radii 0 0 between 50 A and 70 A. Clearly migration of copper from the catalyst pores took place during drying and calcination, as such pores could not enclose the observed crystallite sizes. There­ fore a significant proportion of the copper loading is not effectively utilized. Total surface area however markedly decreased as copper loading increased, indicating that the cata­ lyst pores were being filled to some extent by well-dispersed copper not visible by XRD. When the pore size distributions of catalysts with copper loadings of 15.6 and 4.4 weight percent 139.

were compared with that of the support, no change in the distribu­ tion was noted (Figure 7.1) indicating that pore entrances were not simply being blocked by increased copper loadings.

100 fk • a: .. • 80 en f) _.I z a: 1-- 60 .. SILICA .. 0 t- I LL 40 0 z LCJW CU. .. X z 20 C!) HIGH CU .. C!) z f)

J[• 0 0 20 40 60 80 100 PCJRE RADIUS CANGSTRCJMSl

FIGURE 7.1. Pore size distributions of copper/silica catalysts

It would seem therefore that impregnation of the silica support yields both well dispersed copper residing within the catalyst pores and comparatively crystalline copper external to the pores. In Table 7.1 dispersion has been calculated via the method of Appendix VII for each of the catalysts. Excluding the 140.

catalysts with lowest metal surface areas which would be expected to have very low activity, a copper loading of 15.6 weight percent appears to most effectively utilize the porous nature of the support. In order to observe if copper dispersion affected the rate of methanol dehydrogenation in the absence of support inter­ action, the catalysts with copper loadings of 15.6, 9.1 and 4.4 weight percent were included in the catalyst screening experiments. The characteristics of these catalysts are sununarized in Table 7.2. Hereafter, they will be referred to as having "high, medium and low" copper loadings respectively.

Acatalyst of 20 weight percent copper supported on silica was also prepared using a support of surface area 50 m2g-1• Metal surface area (0.5 m2g-1) and activity were very low indicating a

0 poor degree of copper dispersion. A crystallite size of 400 A reinforced this view and the catalyst was not included in the screening experiments.

7.2.2.2. Ion exchanged catalysts

Exchange reactions between metal cations and catalytic supports produces a uniform distribution of metal ions on the surface of the support, and hence a more highly dispersed catalyst after calcination and reduction. Darling et al. [233] and Benesi et al. [234] have shown-how exchange between Pt(NH3):+ and surface hydroxyl groups of silica gel can produce supported platinum

0 catalysts with crystallite sizes as low as 15 A. Complex ion exchange has also been used by Morikawa et al. [235] to prepare highly dispersed palladium. 141.

Highly active copper catalysts for the steam reforming of methanol (Section 2.3.3.) have been prepared by exchange between hydroxyl hydrogen on the surface of silica supports and copper tetrammine (Cu(NH3)!+). Kobayashi et al. [141,142] have shown how turnover frequency of the steam reforming reaction decreases at high levels of copper dispersion and have proposed that highly dispersed copper was not easily reduced to the active metal, a view supported by Voge and Atkins [195]. Takezawa et al. [214] have used XPS to show that isolated cu 2+, which can remain undecomposed at 500°C, reduces to inactive Cu+. A detailed study of the effect of copper loading and calcination temperature on copper dispersion has been carried out by Shimokawabe et al. [225]. Cu(NH3)!+ was shown to be held on the surface in both isolated and clustered forms, the latter form yielding highly dispersed CuO when calcined at 500°C. The isolated tetrallllline ions, which were preferably formed at low copper loadings (~ 2 percent), formed divalent copper at this calcination temperature.

The preparation of a very highly dispersed copper catalyst for the screening experiments was based on the method of Kobayashi et al. [141]. A solution of copper tetrammine was prepared by slow addition of concentrated ammonia (28 percent) to the cupric nitrate stock solution of Section 4.2.2. until a solution pH of 11 was obtained. The complex cation forms by the reaction

{7.1)

Silica particles {400-600 micron), that had been thoroughly dried and evacuated, were added to the solution and the mixture gently 142. agitated for 6 hours. Ammonia solution was added as necessary to maintain a pH of 11 as required to avoid incorporation of anions into the catalyst [139]. After filtering the silica from the solution, the catalyst was prepared by drying the particles for 16 hours at 110°C followed by calcination at 500°C for 4 hours. The presence of isolated cu 2+ gave the material a distinct green colouring after calcination.

The first exchange catalyst was prepared from a 0.04 M solu- 2+ 2 -1 tion of Cu(NH3)4 and had a metal surface area of 2.3 mg • This was not considered particularly satisfactory as metal surface areas over twice as large were obtained by standard impregnation techni­ ques as shown in the previous section. Hirose et al. [142] have shown that activity for methanol steam reforming could be increased by calcination at temperatures in excess of 700°C, when isolated cu 2+ was converted to CuO. Accordingly the catalyst was re-calcined at 800°C for 3 hours, but copper surface area fell dramatically to 0.08 mg2 -1 •

A second catalyst was_ prepared from a 0.21 M solution of

Cu(NH3)~+ and a considerably higher copper surface area (6.5 m2g-1) was obtained. Further increases in solution concentration yielded only marginally higher copper surface areas, so it was the second I • catalyst that was chosen for the screen1ng experiments. Copper content was found by Atomic Absorption Spectroscopy (Section 4.3.1.) to be 5.0 weight percent. Shimokawabe et al. [225] have indicated that calcination at 500°C of a catalyst of this loading yields isolated cu 2+ as well as highly dispersed CuO. Free CuO was determined by TGA (Section 4.3.4.) to be approximately 3.5 weight 143. percent indicating that isolated cu2+ made a significant contribu­ tion to the total copper loading. Given that metallic copper forms by reduction of CuO, a dispersion of 0.29 can be calculated via the procedure in Appendix VII. Comparison with the results for impregnated catalysts (Table 7.1) shows that a considerably higher degree of dispersion can be obtained using the ion exchange tech­ nique. The results show that for a similar metal loading a dispersion of around 0.03 was obtained for the impregnated catalyst, a value almost one-tenth of that obtained by ion exchange. Further­ more, as compared with the copper powder catalyst of Section 7.2.1., a three hundred-fold increase in dispersion is obtained via the ion exchange technique.

Characteristics of the ion exchange catalyst are sununarized in Table 7.2. Crystallite size was too low to be accurately

0 determined by XRD line broadening, but was certainly less than 30 A.

7.2.3. Copper supported on alumina

Catalysts of copper supported on alumina have been shown by Inui et al. [143] to be effective for the steam reforming of methanol (Section 2.3.3.). Despite the acidity of the support, high selectivities were obtained as adsorption of water on acidic sites prevented the dehydration of methanol by the reaction

(7.2)

The preparation of copper/alumina catalysts has been studied in detail by Friedman et al. [223]. At low copper loadings, cupric oxide was extremely well dispersed and CuO crystallites did 144.

not form below a loading of 4 weight percent. At high calcination temperatures (> 600°C), bulk copper aluminate (CuA1 2o4) was detected by XRD. The deactivation of copper/alumina catalysts at high temperatures has been attributed by Jacobson and Selwood to formation of CuA1 2o4 [236]. Lo Jacono et al. [224] have shown that, although thermodynamically favoured, the reduction of CuA1 2o4 to active Cu(O) is subject to strong kinetic interference. The formation of CuA1 2o4 is minimized if calcination temperature is kept below 600°C [223]. Voge and Atkins [195] have shown that copper/alumina catalysts formed by calcination at 450°C are completely reduced by hydrogen at 200°C.

Hassan et al. [206,207] have studied the effect of copper loading and pore structure on dispersion and activity of copper/ alumina catalysts for decompositionofH2o2 and hydrogenation of maleic acid. Dispersion and activity were found to decrease as copper loading increased.

A copper/alumina catalyst was required for the catalyst screening experiments to examine the effect of an acidic support on the activity of copper for methanol dehydrogenation to methyl formate. The incipient wetness technique of Section 4.2.2. was used to prepare a very effectively dispersed copper catalyst from 2 -1 a y alumina support of surface area 200 mg (Norton Pty. Ltd.). Calcination temperature was kept low (350°C) to prevent the formation of copper aluminate. The results of catalyst charac­ terization by the techniques of Section 4.3 were:- 145.

Copper Content = 11.1 weight percent Total Surface Area = 139 m2g-l Copper Surface Area= 14.0 mg2 -1 0 Crystallite Size = 48 A Dispersion = 0.19

No bulk CuA1 2o4 was detected by XRD.

7.2.4. Copper supported on magnesia

A catalyst of copper supported on magnesia was prepared to examine the effect of a basic support on the dehydrogenation of methanol to methyl formate. A similar catalyst has been used by Hayashi et al. [237] to study the effect of support basicity on the hydrolysis of nitriles to amides.

Catalyst preparation was based on the method employed by Voge and Atkins [195]. However, whereas these authors introduced copper by co-precipitation, the impregnation technique of Section

4.2.2 has been used here. Magnesium carbonate (MgC03) was prepared by the addition of a sodium carbonate (Na2co3) solution to a solution of magnesium nitrate (Mg(N03)2), and after thorough washing, was dried at 110°C for 16 hours. Calcination at 400°C for 16 hours yielded a magnesia (MgO) support of surface area 70 m2g-1• Despite the high copper loading employed in the impreg­ nation procedure, reduction of copper/magnesia in hydrogen yielded a catalyst of relatively high metal surface area and low copper crystallite size, as shown in Table 7.2. Hayashi et al. [237] have also corrrnented on the unusually high dispersion of copper supported on magnesia. Christiansen and Huffman [238] have shown 146. that copper/magnesia catalysts reduce completely in hydrogen at 185°C. Compounds or solid solutions are not known to form below 630°C [195].

7.2.5. Copper supported on chromia

To substantiate the interaction between copper and chromia that was apparent for the copper chromite catalysts of Chapter 6, a catalyst of copper supported on chromia was prepared. The effect of chromia to reduce activity and selectivity for methanol dehydrogenation to methyl formate, should be highly prevalent if the support is 100 percent chromia.

Chromia (Cr2o3) is usually prepared by decomposition of chromium III hydroxide (Cr(OH) 3) under vacuum [239] or under a hydrogen atmosphere [240] at ea. 400°C. Shafer and Roy [241] have shown the transition to occur via the formation of CrO(OH). Chromia was prepared for the catalyst screening experiments by the method of Ciapetta and Plank [240].

Cr(OH) 3 was obtained by addition of 1000 cm3 of a 3 percent NH3 solution to a vigorously stirred solution consisting of 100 g of Cr(N03)3.9H20 dissolved in 1000 cm3 of distilled water and maintained at 60°C. After two hours, the fine precipitate was filtered and dried at 110°C for 16 hours. The hard granules of Cr(OH)3 were then crushed and loaded to a U-tube reactor immersed in a molten salt bath. Hydrogen was passed through the reactor and the bath temperature raised to 400°C. After 6 hours, the chromia was removed from the reactor and crushed and sieved to 400-600 micron particles. Surface area of this support was 110 m2g-1• 147.

TABLE 7.2. Properties of reduced catalyst samples

Catalyst Surface Area Copper {m-29-1) Crystal 1 ite Content Diameter Code Type {wt.%) SBET Seu

A Harshaw 1808 35.5 28.4 6.8 60 copper chromite B Harshaw 0203 59.4 16.3 10.8 100 copper chromite C Girdler G-22* 33.7 45.1 15.4 90 copper chromite D Copper on chromia 5.5 83.1 4.0 166 E Copper on magnesia 20.7 44.9 12.3 50

F Copper on silica 15.6 227 4.8 230 high loading G Copper on silica 9.1 240 1.6 180 moderate loading H Copper on silica 4.4 261 0.94 190 low loading

I Copper on silica 5.0 233 6.5 ion exchange

J Copper oxide 99.0 0.93 280 powder

K Raney copper · 99.3 18.5 18.1 100

* contains 10% BaO 148.

After impregnation with the cupric nitrate stock solution of Section 4.2.2, the catalyst was dried at 110°C for 16 hours and reduced in the reactor under hydrogen for a further 16 hours. As the decomposition temperature .of Cu{N03)2 is 170°C, complete conversion to copper metal would be expected. No decomposition products could be detected by gas chromatography after 16 hours. Calcination was not carried out as the chromia support underwent a dramatic decrease in surface area when heated to 400°C in air. The XRD spectra of the calcined support showed the distinct pattern of chromia indicating that crystallization was the likely cause of the drop in surface area.

As can be seen from Table 7.2, a reasonable copper surface area {4 m2g-1) can be obtained using a 5.5 weight percent copper loading in the impregnation procedure, and it was this catalyst that was used in the screening experiments. Dispersion {0.11) is higher for this catalyst than for the copper chromite catalysts {section 6.4) and the effect of chromia on activity and selectivity should be readily apparent.

7.3. Preliminary Investigations

Initial screening of catalysts involved measurement of conversion, selectivity and stability at 220°C and an LHSV of 26 h-1• Under these conditions, zinc chromite, Pd/Al 2o3 and Ag0/Al 2o3 {Table 4.7) showed no activity at all. If temperature was raised to 300°C. 0.01 percent methyl fonnate was produced over zinc chromite, along with 3 percent carbon monoxide. The other two catalysts produced only traces of carbon monoxide{~ 1 p~rcent) 149. at this temperature. The Cu/Ni0/Al 2o3 hydrogenation catalyst of Table 4.7 produced only 0.26 percent methyl formate plus 0.64 per­ cent carbon monoxide at 220°C. At 300°C, levels of methyl formate and carbon monoxide were 2.17 percent and 4.74 percent respectively. Traces of formaldehyde (ea. 1 percent) were also produced at this temperature. The low selectivity of this catalyst can be attribu­ ted to the activity of nickel .for methanol decomposition by the reaction

(2.31)

A Raney nickel catalyst, prepared from a 48 percent aluminium/SO percent nickel alloy (Davison Chemical Division of W.R. Grace and Co.) by the extraction procedure of Section 4.2.3 gave 10 percent conversion of methanol to carbon monoxide and hydrogen at 220°C. The activity of nickel for methanol decomposition has been explained via dissociative adsorption of methanol on the metal surface [122]. Little could be gained from further experimentation with these low activity catalysts, and so they were not investigated any further.

The Cu/Zn0/A1 2o3 methanol synthesis catalyst of Table 4.7 showed intermediate activity for methanol dehydrogenation to methyl formate with an initial conversion of 14 percent being obtained at 220°C. The level of activity, however, is signifi­ cantly- less than that ,for copper chromite catalysts under equivalent conditions (Figure 6.6). This result is surprising as a very high copper surface area (43 m2g-1) was obtained via the N2o reaction technique described in Section 4.3.2.2. Further, conversion fell to only 6 percent after one hour and initial activity could only be restored by treatment with 1 percent o2 in 150. helium at 300°C. Possibly the incorporation of zinc oxide has a very adverse effect on catalyst activity and stability. As the function of zinc oxide in methanol synthesis is not properly understood, it is certainly beyond the scope of this thesis to examine the effect of zinc oxide in methanol dehydrogenation, especially as the catalyst is not particularly effective in any case. Accordingly, this catalyst will not be considered further. Suffice to say that the claims in Patent literature [6] as to the effectiveness of the catalyst are questionable.

The catalyst of copper supported alumina was tested at 220°C over a range of flowrates. Results are presented in graphical form in Figure 7.2. High levels of dimethyl ether suggest that dehydration activity of the catalyst is very large. Yield of methyl formate is small in comparison. The similarity of carbon dioxide and dimethyl ether levels suggests that conversion to methyl formate may be quite high, though yield is reduced by the following sequence of reactions

2CH30H + HCOOCH3 + 2H2 (1.2)

2CH30H + CH30CH3 + H2o (7.2)

HCOOCH3 + H2o + HCOOH + CH30H (2.53)

HCOOH + Co2·+ H2 (2.54)

Reactions 2.53 and 2.54 would be expected to occur readily (Section 2~3.3) and the overall effect of the reaction sequence would be to produce roughly equal amounts of carbon dioxide and 151.

.2

4> C02

C!l CH30CH3

• HCOOCH3 .15

• 1 z ...... l!:J t- u a: a::: LL w _J .05 l!:J L

'

...... 0 0 20 40 60 80 100 120 LHSV ( /HRJ

FIGURE 7.2. Effect of LHSV on product distribution for copper alumina. 220°( 152. dimethyl ether. Small quantities(< 1.5 percent) of a higher molecular weight species were also detected. Although no stan­ dards were available, the formation of methylal (CH30cH20cH3) by the reaction

(7.3) is suspected in light of the results of Ai [129] who has observed the formation of methylal in the oxidation of methanol to methyl formate over catalysts with acidic functions.

Clearly the effect of support acidity is to lower catalyst selectivity by promotion of methanol dehydration. Due to the very low selectivity(< 3 percent) obtained under the test conditions, this catalyst wa5 not examined any further.

The activity and selectivity of the remaining copper cata­ lysts will be discussed in Section 7.4.

7.4. Results

The copper catalysts of Table 7.2 were active in varying degrees for the dehydrogenation of methanol to methyl fonnate. Catalysts A, C, E, F, G, Hand I exhibited stable activity over a period of three hours. However some deactivation was evident after longer periods on line. The mild deactivation of catalyst B has been discussed in Section 6.4. Some deactivation was evident for catalysts D and J, as shown in Figure 7.3. Since initial conversions were low, initial rate of reaction was obtained by extrapolation to conversion at time equal to zero. 153.

10

9

8

7

6 ..... C0PP£R POVD£R -N z 5 0 -(/) et= 4 l1J > z 0 u 3

CfJPPER/CHROMJA 2

l

0 0 10 20 30 40 50 60 TIME (MINUTES)

FIGURE 7.3. Deactivation of copper powder and copper chromia. 220°C. LHSV = 26 h-1 154.

Excluding catalyst J, all catalysts were tested for activity and selectivity over the range of flowrates 14 h-l < LHSV < 200 h-1• Results are tabulated in Appendix VIII, and presented in graphical fonn in Figures 7.4 to 7.10. High activity was obtained for the copper chromites and catalysts E, F and I, reflecting the high copper surface areas of these catalysts. No products other than hydrogen, methyl fonnate, carbon monoxide and carbon dioxide were detected, though carbon dioxide levels never exceeded 0.5 percent. Selectivities in excess of 90 percent can be obtained for all catalysts except catalyst D. In Figure 7.11, the data of Appendix VIII have been used to illustrate the relationship between selec­ tivity and conversion.

To determine the effect of reactant concentration, the feed of methanol was maintained at a constant LHSV (53 h-1) and conver­ sion was measured at various levels of helium dilution. In Figure 7.12, conversion is plotted against methanol mole fraction for five of the catalysts. Clearly, in the range of methanol concen­ trations encompassed in Appendix VIII, the rate of reaction is independent of methanol concentration indicating that zero order kinetics are being obeyed.

The chromia, magnesia and silica supports were totally inactive for methanol dehydrogenation at 220°C.

7.5. Discussion

The results of sections 7.3 and 7.4 show that copper cata­ lysts are active and selective for the dehydrogenation of methanol to methyl formate, and that under the conditions of experimentation, 155.

0 0 N

Ln "-

0 -Ln

Ln -N

0 -0 -a:: ::r: Ln " -' > Cl) 0 ::c Ln _j

Ln N

0 ,..,0 Ln 0 Ln 0 N N - - Ln 0 (1.l NrJ1SH3ANrJJ

FIGURE 7.4. Effect of LHSV on convers;on for copper chromite catalysts at 22n°c 156.

a a N

l/) "-

a l/)-

l/) -N .,., N CD N 0 0 CD N CJ 0 a - 0 ~ El • - ,..... a::: ::r: l/) " -' > (/) a ::r: l/) _J

u:, N

a 0 a a 0 0 0 0 - 0) CD " U) l/) (t.l .l1IAI1J3l3S

FIGURE 7.5. Effect of LHSV on selectivity for copper chromite catalysts at 220°c 157.

0 0 N

LI) "-

0 LI)- :a: :::>..... ::c LI) C, C .... w N ::c :a: - .. H

a: 0 ....u 0 ...... J - U) ~ :::> ~ 'u I LI) " '- > (/) 0 I LI) _J

• LI) N

0

LI) 0 LI) 0 0 LI) 0 ,.,, N N - - C i. J NCI SH3ANCJ

FIGURE 7.6. Effect of LHSV on conversion for copper/ silica catalysts at 220°c 158.

D D N

-

D -Ln ::c ::, :x:: CJ -Cl s Ln llJ 0 -:x:: ::c --' -N

> Cl) D :r: Ln _J

' Ln N

D D D D D D D D - 0) co f'- <.O Ln (1.l .l1IAI1J3l3S

FIGURE 7.7. Effect of LHSV on selectivity for copper/silica catalysts at 220°c 159.

0 0 EJ N

w Ln C!Iz « « .... " V) - « ::c u u w .... X z C!I ..J LLI « .... 0 V) z X: a Ln ::l .... ::l' u' u - ~ EJ

Ln -N

0 -0 O:'.: I Ln " .....,' > en 0 I Ln _J

Ln N

0

0 Ln 0 0 ,.., ...... 0 N N - [ t. l NCJ I SH3ANCJJ

FIGURE 7.8. Effect of LHSV on conversion for copper/ silica (ion exchange) and copper/magnesia catalysts at 220°c 160.

0 0 N

IJ) r--. "i

w C, z 0 ct ct IJ) ct :I: V) u u -w - ..... X z _J w l!I z ct - X: 0 IJ) .... :::> N ~ ' u - El

0 0- ~ n::: I IJ) r--. -' > en 0 I IJ) _J

IJ) N

C)

C) C) C) C) C) C) 0 IJ) - 0) CD r--. (0 (%) J..1IAI1J3l3S

FIGURE 7.9 Effect of LHSV on selectivity for copper/ silica (ion exchange) and copper/magnesia catalysts at 220°c 161.

60 ...... ;-,-.: ...... , 55 >- t- -> -t- 50 u w -' w (/) 45

40 0 25 50 75 100 LHSV /HRl

10

...... ;-,-.: ...... , 7.5 z ~ -(/) a::: 5 w > z ~ u 2.5

0 0 25 50 75 100 LHSV ( /HRl

FIGURE 7.10. Effect of LHSV on selectivity and conversion for copper/chromia catalyst at 220°c · 162.

0 tt)

l/) N

a: cn a: ,...... -w 0 -:c N ~ z D CJ D:: '-J a: :I: CD :c u 0 z CD :::) :::) ..... u' 'u 0 w l/) C, z -Cl) a: • • • - a:::: :I: u w >< w > :I: z C, z 0 0 D u -:I: .... - a: a: ....u .... u ...... J ...... J I') cn en N l/) 0 N N :::) :::) 0 'u ' u c.:,

EJ ... M E)

0 0 0 0 0 0 0 0 0 - en CD " (.0 l/) od" ( 1/. ) J..11 t\ I 1J:3T3S

FIGURE 7.11. The relationships between conversion and selectivity for copper catalysts at 220°C 163.

20

18

16

)

14

12 10 m m ... - [!) 10 ;-..: ...... , z 8 D V -en l( ~ w 6 (!) CU/SILICA (HIGH) > z D 0203 u • 4 I!) 1808

A G 22

2 X CU/MAGNESIA

0 1 .9 .8 .7 .6 .5 .4 M~LE FRACTION OF METHANOL IN FEED

FIGURE 7.12. Effect of methanol concentration on conversion for copper catalysts. 220°C, LHSV = 53 h-1 164.

zero order kinetics are obeyed. The latter result is not surpri­ sing as complete surface coverage with methanol would be expected at the high inlet concentrations employed. In Table 7.3, the kinetic expression of Miyazaki and Yasumori [122] {Equation 2.37) has been used to predict the effect of methanol concentration on reaction rate at 220°C, assuming a "roughness factor" of unity. Clearly, methanol concentration does not appreciably influence rate above a mole fraction of 0.1.

TABLE 7.3. Effect of methanol concentration on reaction rate as predicted from the kinetic expression of Miyazaki and Yasumori [122] {Equation 2.37)

Mole Fraction of Methanol Rate x 10-16 {molecules cm-2sec-1)

1.0 1.73 0.8 1.73 0.6 1.72 0.4 1.71 0.2 1.66 0.1 1.57 0.05 1.42 0.01 0.81

In the case of a zero order chemical reaction, rate of reac­ tion cannot be limited by pore or bulk diffusion. Rate control by product desorption is unlikely as the heat of adsorption of methyl formate on copper, found by Miyazaki and Yasumori [122] to be approximately 30 kJ mol-1, is low compared with the value of 165.

80 kJ mol-l calculated in Section 6.4 for the activation energy of the dehydrogenation reaction on copper chromite catalysts. In any case, a particularly strong form of bonding between methyl formate and copper is hard to envisage. It is proposed then, that the rate limiting step in the dehydrogenation of methanol to methyl formate is thesurface reaction of methanol, in agreement with the conclu­ sions of Miyazaki and Yasumori [122].

Under the conditions of zero order kinetics, rate and rate constant are equivalent, and hence rate constant can be obtained directly from the data of Appendix VIII by the procedure of Section 4.1.4. As the chromia, magnesia and silica supports were totally inactive for methanol dehydrogenation at 220°C, specific rate constants can be calculated using Equation 4.10. The values of rate constant and specific rate constant, hereafter referred to as "specific activity", are listed in Table 7.4.

Specific activity does not vary appreciably between cata­ lysts F, G, Hand J indicating that, in the absence of support interaction, activity is proportional to copper surface area. The specific activity of catalyst I is slightly lower by comparison, probably due to support interaction of the type described by Kobayashi et al. [141,142] for the steam reforming of methanol over copper/silica catalysts. Although sili~a is normally consi­ dered to be an inert support, it was shown that activity for steam reforming was adversely affected when metal surface area exceeded a value of 50 m2 (gram copper)-1. As the metal surface area of catalyst I is equivalent to 130 m2 (gram copper)-1, a degree of support interaction would not be unexpected. 166.

TABLE 7.4. Kinetic parameters for the dehydrogenation of methanol

Catalyst Rate Constant Specific Rate Constant x105 xl05 (mole s-lg-1) (mole s-lm-2)

A. Harshaw 1808 Copper Chromite 5.84 0.86 B. Harshaw 0203 Copper Chromite 6.73 0.62 c. Girdler G-22 Copper Chromite 6.03 0.39 D. Copper on Chromia 0.89· 0.22

E. Copper on Magnesia 8.24 0.67 F. Copper on Silica High Loading 11.04 2.28 G. Copper on Silica High Loading 4.32 2.77 H. Copper on Silica Low Loading 2.24 2.38 I. Copper on Silica Ion Exchange 11.58 1.78 J. Copper Powder 2.18 2.40

When the threehundred-fold variation in dispersion amongst the copper/silica and copper powder catalysts (Section 7.2.2.2} is taken into account, the variations in specific activity are trivial. If particular irregularities within the copper crystallite struc­ ture were important to catalytic activity, it would follow that catalyst I, with the highest degree of dispersion, should exhibit a markedly higher specific activity than the other copper/silica catalysts. Since this is not the case, the dehydrogenation 167.

reaction can be described as "facile", i.e. reaction rate depends simply on the number of surface copper atoms available, and not on the number of corner and/or edge atoms [230].

However, factors other than metal surface area appear to be significant in determining the activity of the chromium-containing catalysts. In Table 7.4, the average specific activity of the copper/silica catalysts is between three and ten times higher than the specific activities of catalysts A, B, C and D. A number of explanations can be put forward to account for the low specific activity of these catalysts.

Firstly, specific activities could be made artificially low by an overestimation of metal surface area due to bulk oxidation or decomposition of nitrous oxide by the support material. Bulk oxidation has already been shown to be insignificant in Section 5.2.3. Further, neither cupric chromite, chromia, nor a mixture of chromia and cuprous chromite prepared by the technique of Miya et al. (175], showed any activity for the decomposition of nitrous oxide. Clearly metal surface areas of the chromium-containing catalysts have not been overestimated, and real values for speci­ fic activity are being observed.

Secondly, specific activity of the copper chromite catalysts could be decreased by pore diffusion control. Although the pore size distributions of two copper chromite catalysts (Figure 4.1) show that a significant degree of the surface area is contained 0 within relatively small pores(< 50 A), it has already been stated that zero order kinetics could not be observed if resistance due 168.

to pore diffusion was a contributing factor in the observed reac­ tion rates. Further, as the combination of zero order kinetics and rate control by surface reaction implies saturation of the surface by adsorbed methanol, a reduction in rate due to an excess of adsorbed hydrogen or methyl fonnate, is not a feasible explana- .tion, especially as the differences in specific activity were evident at the lowest levels of conversion. Hydrogen is likely to be adsorbed on chromi.a [186,187], and the results of Higdon et al. [44] as to decarbonylation activity of chromia, suggest methyl fonnate is also likely to be adsorbed. However, as the reaction is entirely specific to the availability of copper surface area, adsorption of products on chromia can be ignored.

The most plausible explanation for the low specific activi­ ties of the chromium-containing catalysts is an electronic interaction of the type discussed in Chapter 6 to explain diffe­ rences in activity amongst the copper chromite catalysts. The unexpectedly low activity of catalyst C was explained by proposing an interaction between copper and chromia that served to negate the effect of the very high metal surface area of that catalyst (Table 5.1). It can be seen from Table 7.4 that, of the three copper chromite catalysts tested in Chapter 6, catalyst C shows the lowest specific activity. The effect of chromia is confirmed by the particularly low specific activity of catalyst D. However, catalysts A and B, which contained 111.1ch lower levels of chromia (Table 5.8), also have low specific activities when compared with the copper/silica catalysts indicating that the interaction is not specific to one particular type of chromium III oxide. In the 169. case of catalyst B for instance, the interaction could only arise from the presence of cuprous chromite.

Although the strength of the interaction between copper and chromia is readily apparent from Table 7.4, the nature of the effect is difficult to describe. Tauster et al. [242] have described how catalyst reduction can lead to electron transfer from a partially reducible support, such as titania (Ti02), to the active metal. A change in the electronic configuration of the metal atoms can result in a modification of the catalytic activity. Niiyama et al. [243] have noted that copper/titania catalysts are much less active than copper/silica catalysts for methanol dehydro­ genation. Recently, Severino and Laine [168] have shown the activity of copper chromite catalysts for oxidation of carbon monoxide to be affected by electron transfer between chromia and copper. Selwood et al. [171] used measurements of magnetic suscep­ tibility to identify electronic interactions within the copper chromite system.

A low specific activity is also evident in Table 7.4 for copper supported on magnesia (catalyst E). Once more, mass trans­ fer constrictions and overestimation of copper surface area can be ignored, and electronic interaction is again suggested, especially ' in light of the low copper crystallite size (Table 7.2). Although magnesia is not normally considered to be a reducible oxide,Dowden [244] has discussed the weak chemisorption of hydrogen on magnesia in tenns of localized charge transfer. In Section 6.4 the effect of barium oxide in catalyst Chas been discussed with refe~ence to the low specific activity of the copper/magnesia catalyst. 170.

Figure 7.11 shows the differences in selectivity encountered amongstthecopper catalysts. The low selectivities of catalysts D and E were suspected of being related to the activity of chromia and magnesia for the decarbonylation of methyl formate by reaction 2.12 [44]. To test this theory, a 2:1 mixture of nitrogen and methyl fonnate was passed over silica, magnesia and chromia supports at 220°C and as Table 7.5 shows, the activity of the supports for the decarbonylation reaction reflected the selectivity differences evident in Figure 7.11. A particular activity of copper for decarbonylation is evident from the tendency of catalysts F and I towards low selectivity at high conversions to methyl formate, and is consistent with the findings of Schwab and Knoezinger [193]. As both copper and chromia possess decarbonylation activity, lower selectivities than were obtained might have been expected for catalysts A and C, which have been shown to consist of copper supported on chromia after reduction (Section 5.3).

TABLE 7.5. Decarbonylation activity of support materials 30 cm3min- 1 HCOOCH3 at 220°C Support Mass_-{g) Conversion by Equation 2.12 (%)

Silica 0.5 0 Magnesia 0.25 0.7 Chromia 0.5 8.0

Given that total surface areas in Table 7.2 reflect the availabi­ lity of decarbonylation sites, the low total surface area of catalysts A and C, as compared with catalyst D, is probably an 171.

important factor in maintaining selectivity of the copper chromite catalysts at acceptable levels.

An interesting result is the particularly high selectivity of catalyst B. This catalyst contains no chromia after reduction (Table 5.8), and it is interesting to note that selectivity is higher for this catalyst than for catalysts F and I. Some form of interaction between copper and cuprous chromite was suspected, but not further investigated. The relationship between the high selectivity and tendency of catalyst B towards deactivation has been discussed in Section 6.4.

7.6. Catalyst Selection

The rates of methanol dehydrogenation on the most active catalysts (A, B, C, E, F and I) are compared on the basis of unit volume of catalyst bed in Table 7.6. This comparison, which takes into account bulk density differences between the catalysts, provides a suitable basis for selecting the most desirable catalyst from an industrial viewpoint.

Clearly, the low specific activities of the copper chromite catalysts (Table 7.4) are more than outweighed by their high bulk densities. Catalyst B, the Harshaw 0203 copper chromite, emerges as the most active catalyst, but, as shown in Section 6.3.2, 1s prone to deactivation. The most desirable catalysts appear to be catalyst A, the Harshaw 1808 copper chromite, and catalyst C, the Girdler G-22 copper chromite, which combine high activity and long active life. As shown in Figure 6.11 relatively high selectivities can be obtained with either of these catalysts. 172.

TABLE 7.6. Rates of dehydrogenation on the more active catalysts

Catalyst Rate (kg methanol hour-1(litre catalyst)-1)

A. Harshaw 1808 Copper Chromite 9.96 B. Harshaw 0203 Copper Chromite 15.04 c. Girdler G-22 Copper Chromite 8. 78

E. Copper on Magnesia 4.10 F. Copper on Silica High Loading 5.47 I. Copper on Silica Ion exchange 5.48

7.7. Conclusion

The activity of a range of metal and metal oxide catalysts for the dehydrogenation of methanol to methyl formate has been investigated. Copper catalysts are by far the most active and selective for the reaction, and in general, high activity is favoured by high copper surface area. Due to the high concentra­ tions of methanol employed, zero order kinetics were obeyed with the surface reaction of methanol being rate controlling.

Differences in activity and selectivity amongst the copper catalysts can be related to the influence of support material. When an inert silica support was used, activity was solely depen­ dent on available copper surface area. Highest specific activities were obtained for such catalysts. Specific activities of the copper chromite catalysts were low in comparison and consistent with the idea of electronic interaction between chromium III oxides and copper. The relatively low selectivities of chromia and magnesia- 173.

supported-catalysts can be related to the activity of the supports for methanol decarbonylation. Methanol dehydration is promoted by the use of an acidic alumina support.

When bulk density is taken into consideration, the copper chromite catalysts emerge as being most favourable for the dehydrogenation reaction. 174.

CHAPTER 8

BEHAVIOUR OF RANEY COPPER CATALYSTS

8.1. Introduction

Comparison of copper catalysts on the basis of specific activity as carried out in the previous chapter, showed the most active catalysts for methanol dehydrogenation to methyl formate to be pure copper powder, or copper supported on an inert carrier (silica). However the practical activity of these catalysts was limited by the low copper surface area of the powder and low bulk density of the silica-supported catalysts. It was hoped that a Raney copper catalyst would overcome these disadvantages.

As has been shown in Section 4.2.3, leaching of copper/ aluminium alloys yields catalysts with copper surface areas as high as 18 m2g-1, and since the leaching process can be made virtu­ ally quantitative to yield almost pure skeletal copper, no support interaction could be envisaged. Furthennore, if the specific activities of catalysts F, G, Hand Jin Table 7.4 are used to predict an average specific activity for copper of 2.Sxlo-5 mole s-lm-2, then given that the bulk density of Raney copper is approximately 1 g cm-3, a rate per unit volume of catalyst bed of up to 50 kg methanol h-l ( l catalyst)-l can be calculated for the Raney copper catalysts of Table 8.1. This compares very favour­ ably with the equivalent rates for the more active supported catalysts (Table 7.6). Selectivities of the Raney copper catalysts would be expected to be very high due to the absence of support

(J"I (J"I

......

. .

"' "'

/g) /g)

2.6 2.6

2

1.2 1.2

3.9 3.9

8.6 8.6

18.1 18.1

31 31

(m

Copper Copper

Areas Areas

12.0 12.0 15 15

21.2 21.2

18.6 18.6

19.9 19.9

45 45

Total Total

Surface Surface

Properties Properties

Catalyst Catalyst

Content Content

79.6 79.6

95.6 95.6

63.8 63.8

99.3 99.3

{wt%) {wt%)

Copper Copper

3 3

3 3

0.5 0.5

3 3

3 3

62. 62.

{h) {h)

Time Time

Leaching Leaching

catalysts. catalysts.

10 10

NaOH NaOH

40 40

40 40 40 40

40 40

40 40

(wt%) (wt%)

Ltd.) Ltd.)

Concentration Concentration

copper copper

-

Pty. Pty.

Raney Raney

10 10

Cr Cr

of of

hours hours

Catoleum Catoleum

3 3

by by

15 15

Zn Zn

after after

Characteristics Characteristics

Composition Composition

NaOH NaOH

{Supplied {Supplied

{wt%) {wt%)

50 50

50 50

48 48

50 50

Al Al

2 2

8.1. 8.1.

Fresh Fresh

Alloy Alloy

CuA1

1. 1.

2. 2.

50 50

50 50

50 50 50

Cu Cu

35 35

41 41 49

TABLE TABLE ls2 ls2 176. material that might promote the decarbonylation of methyl formate (Equation 2.12).

Charles and Robinet[lll.245] have patented a Raney copper catalyst for the dehydrogenation of methanol to methyl formate at 180-230°C. It was claimed that the catalyst was more active. more selective and exhibited a longer active life than conventional supported copper catalysts. Similar claims have been made for the dehydrogenation of ethanol on Raney copper [246].

The test conditions established in the previous two chapters were used to investigate the effectiveness of Raney copper catalysts for the dehydrogenation reaction. Unless otherwise stated. the catalyst used in the experiments was obtained by leaching a 52 weight percent copper/48 weight percent aluminium alloy over a period of 6 hours. Characteristics of each of the Raney copper catalysts are summarized in Table 8.1.

8.2. Results

The activity of Raney copper for dehydrogenation of methanol to methyl formate was tested at 220°C and an LHSV of 26 h-1 • As can be seen in Figure 8.1. although initial methanol conversion was very high and approached the equilibrium value of 40 percent (Figure 6.1), rapid deactivation occurred. After only thirty minutes, conversion had fallen to level off at a residual value of approximately 10 percent, which compared unfavourably with the activity of most of the copper catalysts already tested (Section 7.4). Selectivity rose slightly during the course of the deacti­ vation since carbon monoxide formation by methyl formate 177.

40

35

30

25

20 '""' X ....., z 15 D ...... U1 a::: w > z 10 D u

5

0 0 10 20 30 40 50 60 TIME (MINUTES)

FIGURE 8.1. Deactivation profile of Raney copper. 220°C, LHSV = 26 h- 1 178. decarbonylation (Equation 2.12) was limited by the progressively lower concentrations of methyl formate. Approximately 95 percent selectivity was obtained after ten minutes on line. However, even at the high levels of conversion obtained after only two minutes on line, selectivity was greater than 85 percent. Traces of formaldehyde, detected over no other copper catalyst, were observed after ten minutes on line. As Figure 8.2 shows, the concentration of formaldehyde increased slightly as deactivation proceeded. Apart from the usual traces of carbon dioxide(< 0.5 percent) no other reaction products were obtained.

The catalyst could be completely reactivated by flushing with hydrogen, helium or nitrogen (100 cm3min- 1) at 220°C over a period of three hours. A reactivation profile, where conversion after two minutes on line has been measured after different periods of flushing with hydrogen, is presented in Figure 8.3. Successive deactivation profiles were always very reproducible.

Given the stability of the copper catalysts tested in the previous chapter, the deactivation of Raney copper was unpreceden­ ted and could not irrmediately be explained. In an attempt to gain more information about the nature of the deactivation, a number of reaction parameters were varied. In Figure 8.4, deactivation profiles have been obtained at half the space velocity and at half the inlet concentration of methanol. The change in space velocity using a pure methanol feed has little effect but the decrease in concentration retards the rate of deactivation and leads to a marked increase in the residual level of conversion. The effect of temperature on the deactivation profile is shown in Figure 8.5. 179 . . OlU D co •r- 0 l/) l,...N :, N "'O

Q) "'O I,.. D >, Q) E) ~ .r:. 0. Q) 0. "'O 0 -(/) .- u It! w E >, D ..,_ I,.. Q) f") => 0 C ~ 0:::"' z ~ ...... 0~ L 0 D C .-4 0 C I N - •r-..., .....0 .r:. w It! ..., .\0 L ...,I- It!> N ...... C•.- II D Cl)..., ..,_ UU> C It! Vl - 0 Cl) :::c u "'O _J . 0 N . co tl) ~ f") N 0 L&J 0::: - ::, C,...... 00001 X (QHJHJJ.. LL..

40 -X z ID ...... 30 (/) a:: w > z 20 ID u

10

0 0 50 100 150 200 TIME (MINUTES)

FIGURE 8. 3. Reactivation profile of Raney copper at 220°C and 100 cmlmin-1 H2. Activity tested at 220°C, LHSV = 26 h-1 180.

40

35 • YCCH30H)•1.LHSV•26/HR

30 X YCCH30H)•1.LHSV•13/HR

25 ~ YCCH30H)•O.S.LHSV•13/HR

,.....,. 20 ::--:...... z ...... D 15 en 0::: w > z 10 D u

5

0 0 10 20 30 40 50 60 TIME CMINUTESl

FIGURE 8.4. Effect of methanol concentration and space velocity on deactivation of Raney copper at 220°c 181.

40

~ 200 C. LHSV= ll/HR

35 ~ 220 C. LHSV=26/HR

• 240 C. LHSV=52/HR 30

25

20 .~ . ·"'...... z a ·15 -(f) a::: w > z 10 a u

5

0 ------t 0 10 20 30 40 50 60 TIME (MINUTES)

FIGURE 8.5. Effect of temperature on deactivation of Raney copper 182.

LHSV was varied so as to achieve similar initial conversions at each temperature. The rate of deactivation is slightly lower at 200°c.

Since variations in reaction conditions appeared to have little effect on the deactivation profile, changes were made to the characteristics of the Raney copper catalyst. Copper surface area was decreased by employing a shorter leaching time, and a lower concentration of sodium hydroxide solution. These changes in the leaching parameters yielded catalysts with copper surface areas of 2.6 m2g-l and 1.2 m2g-l (Table 8.1). As Figure 8.6 shows, the deactivation profiles, and the steady state levels of conver­ sion were not particularly sensitive to large variations in active surface area. Possibly the differences in copper surface area affect the initial activity which, due to the very high rate of deactivation, is not observable. Catalysts with copper surface areas of 18 m2g-l and 9 m2g-l had virtually identical deactivation profiles. In Figure 8.7, deactivation profiles are presented for Raney copper catalysts prepared from copper-aluminium, copper­ chromium-aluminium and copper-zinc-aluminium alloys. The incorporation of catalytic components other than copper does not appear to reduce the extent of deactivation, despite the fact that, in the case of the copper-zinc catalyst, the copper surface area is very high (Table 8.1). It should be added that the unexpectedly low activity of the zinc containing catalyst is consistent with the results obtained in Section 7.3 for the copper-zinc methanol synthesis catalyst. The copper surface area of the chromium­ containing catalysts is of a similar order to the correspondJng 183.

30

~ 4~% NaOH; 3 hours

25

[!) 40% NaOH; 0.5 hours

20

,...... ,. ... 10: ~aOH; 3 hours -X 15 z E) -en 0::: w 10 ... > ~ z E) u

5

0 0 20 40 60 80 100 TIME (MINUTES)

FIGURE 8.6. Deactivation profiles of Raney copper catalysts prepared under different leaching conditions. 220°c, LHSV = 26 h-1 184.

30

25 ~ COPPER

1!1 COPPER-CHROH I UH

20 • COPPER-ZINC

"""" -X 15 z ...... lD (I) 0::: w 10 > z lD u

5

0 0 20 40 60 80 100 TIME (MINUTES)

FIGURE 8.7. Deactivation profiles of Raney catalysts prepared from various alloys. 220°C, LHSV = 26 h-1 185.

values for the Raney copper catalysts, and despite a possible decrease in initial activity, the catalyst behaves essentially like the Raney copper catalysts of Figure 8.6. As the nature of the deactivation was not as yet understood, it was felt that the incorporation of catalytic components other than copper was an unnecessary complication and hence the zinc and chromium-containing Raney catalysts were not investigated further.

The reversibility of the deactivation suggested strong chemisorption of a species present over the catalyst in trace quantities such as carbon monoxide, carbon dioxide or water. To test if chemisorption of carbon monoxide was responsible for deactivation, the methanol feed was diluted 50 percent with a mix­ ture of 40 percent carbon monoxide in helium. As figure 8.8 shows, the rate of deactivation did not decrease with added carbon monoxide but rather was retarded to the degree evident when the methanol feed was diluted 50 percent with helium. Clearly, carbon monoxide was not a poison but was simply acting as a diluent.

The addition of water to the methanol feed produced a 111Jch more marked effect on the deactivation profiles. As shown in Figure 8.9, an increase in water concentration to 5 percent resulted in a marked increase in the residual level of conversion. However, due to the steam reforming mechanism of Section 2.3.3, llllch of the methyl formate reacted further to produce carbon dioxide and hydrogen. Since increased conversions could be achieved in the presence of carbondioxidelevels of up to 10 per­ cent, chemisorption of carbon dioxide could be ruled out. 186.

~0------, l

30

-~ -• z 20 0 V, er w z> 0u ·-----· • 10

QL------+-----....------;---:---- 0 20 4 0 60 80 TIME ON LINE (min.)

FIGURE 8.8. Effect of CO on deactivation of Raney copper. 220°(; LHSV = 13 h- 1; A YCH OH = 1 ; 3 • YCH OH = 0.5; 3 • YCH OH = 0.5, vco = 0.2 3 187.

------, ~

0 •I U) -C E -w z 0 _J ~ z 0 w 11 ~ t-

]I 0 • N /2j ~

L....-t------+------t------o 0 0 0 N - (%) NOIS ~3ANOJ

FIGURE 8.9. Effect of water concentration on deactivation of Raney copper. 220°C, LHSV = 26 h-1 • YH O < 0.002;• YH O = 0.05;.A YH O = 0.01 2 2 2 188.

8.3. Discussion

A colffllon cause of catalyst deactivation is the irreversible adsorption of poisons onto the active surface via compound fonna­ tion between surface metal atoms and the poison. Copper catalysts are known to be easily poisoned by S, Se, Te, P, As, Sb, Bi, Zn, halides, Hg, Pb, NH3, pyridine, c2H2, H2S and PH3 with reactivation being achieved by an oxidation/reduction cycle or by chemical treatment [247]. Inui et al. [143] have attributed the deactiva­ tion of copper/alumina catalysts for the steam reforming of methanol (Section 2.3.3) to oxygen uptake leading to the formation of inactive copper oxides. As Raney copper catalysts could be reactivated simply by flushing with an inert gas, deactivation due to the uptake of oxygen or any other irreversible poison is very unlikely. Furthermore, such deactivation processes would not be expected to affect Raney copper only, and not the other copper catalysts.

The reversibility of the deactivation indicates a weaker form of adsorption. As water is known to be adsorbed onto copper catalysts (101] and was present in trace quantities (ea. 0.2 per­ cent) in the methanol feed, deactivation due to the adsorption of water was considered a likely possibility. Adsorbed water would be readily removed by flushing with an inert gas and so the deactiva­ tion would be reversible. Since water would be expected to react with methyl formate by reaction 2.53 to yield formic acid, deactivation might also arise from the formation of surface fonnate ions (HCOO-) known to be an intermediate in the decomposition of formic acid by the reaction [137] 189.

(2.54)

This form of deactivation would also be reversible as the surface formate ions would eventually decompose [137]. However, as increased water concentrations were shown to actually retard deactivation (Figure 8.9), coverage of the surface by adsorbed water or formate ions cannot be claimed to cause the deactivation of Raney copper catalysts.

The adsorption of carbon monoxide or carbon dioxide has also been ruled out as increases in the concentration of these species did not accelerate deactivation. 'In any case, strong adsorption of carbon monoxide, carbon dioxide, methyl formate or hydrogen is difficult to reconcile with the stability reported in the previous chapter for copper/silica catalysts which should essentially behave as pure copper also.

The complete, but slow reactivation represented in Figure 8.3 suggested desorption of a high molecular weight species, and as traces of formaldehyde were detected in the reactor effluent (Figure 8.2), formaldehyde polymerization seemed a possibility. Although the concentrations of formaldehyde were small, higher surface concentrations are implied, and as formaldehyde is known to readily polymerize to solids of low vapour pressure [190], it is proposed that deactivation is due to the blocking of active surface by the growth of formaldehyde polymers.

A mechanism of deactivation has been envisaged in Figure 8.10. formaldehyde, formed by the surface reaction of adsorbed methanol (r1), can either react to form methyl formate (r2) or undergo CH 3 0 H(g) HCHO(gl r2 l 1~ > HCOOCH3 (gl r1 ~ CH3 OH(al HCHO(al

1l r3 Polymer

...... I.O FIGURE8.10. Mechanism for the deactivation of Raney 0 copper catalysts . 191.

polymerization (r3). As the polymerization process is reversible [190], a small gas phase concentration of formaldehyde would not be unexpected. During the flushing procedure, polymer is broken down by depolymerization to adsorbed formaldehyde which is removed from the catalyst surface by diffusion and by reaction to methyl formate, methanol and tar.bo.n monoxide. It was noted that a steady concentration of carbon monoxide of approximately 0.05 percent could be detected in the reactor effluent over much of the reacti­ vation period and could be equated to desorption of approximately l.3xlo20 atoms of carbon. A Raney copper catalyst with a metal surface area of 18 mg2 -1 conta1ns. 2.63xl0 20 surface atoms of copper, based on a surface density of 1.46xlo19 atoms m- 2 [216]. On the basis of these simple calculations, it is readily seen how the quantity of desorbed material could account for blockage of a large portion of the available copper surface area.

Attempts to measure a decrease in copper surface area on a deactivated Raney copper catalyst were unsuccessful as the prepara­ tion procedure for the nitrous oxide reaction technique of Section 4.3.2.2 flushed away much of the adsorbed material.

As has been mentioned, the reactivation process will involve the depolymerization of formaldehyde polymer and the diffusion of formaldehyde from the catalyst surface, with one of these steps being rate controlling. For the case of catalyst reactivation with hydrogen (Figure 8.3), the mutual diffusivity of fonnaldehyde and hydrogen can be calculated from the relationship [248[

D = 0.001858 312 (1/MA + 1/M )½ r 8 (8.1) p 0 AB nD 192.

where D = mutual diffusivity ( cm 2s -1) T = temperature (K) MA,MB = molecular weights of hydrogen and formaldehyde p = total pressure (atm) 0 Lennard-Jones force constants (A) 0 AB = n = collision integral.

Assuming the diffusion properties of formaldehyde are similar to 0 those of methanol, values for aAB and n0 of 3.227 A and 0.9576 respectively can be determined from tabulated data [248], and a diffusivity of 1.49 cm2s-l can be calculated from Equation 8.1. This value is quite high, and as the quantity of formaldehyde that must be dispersed is undoubtedly small, it is unlikely that the rate of reactivation is governed by fonnaldehyde diffusion. The diffusivities of formaldehyde decomposition products under the conditions of reactivation are also likely to be high and it seems much more reasonable to expect that the rate of reactivation is controlled by the rate of depolymerization of the high molecular weight formaldehyde polymer. Walker [190] and Melia [191] have described the mechanism of the depolymerization process.

If the deactivation of Raney copper catalysts is attributed to polymerization of fonnaldehyde, the insensitivity of the initial stages of deactivation to changes in space velocity, concentration, temperature and copper surface area, is not unex­ pected. Production of formaldehyde by r1 and hence polymerization of formaldehyde by r3, will be at a maximum at the shortest possible time on line when catalyst activity is at its highest. The shapes of the deactivation profiles suggest that the maximum 193. extent of deactivation is virtually being achieved within a couple of minutes on line. The initial rate of deactivation, that would be proportional to space velocity, concentration, temperature and metal surface area, cannot possibly be observed.

More intriguing is the insensitivity of the latter states of deactivation to the same set of reaction parameters. At no stage did a Raney copper catalyst deactivate completely, indicating that a fraction of the copper surface area was inmunefromdeactivation. Yet, as Figure 8.6 shows, the residual activity appears to be largely independent of copper surface area. It is likely that the pore structure of the catalyst is of importance and a brief explanation of catalyst morphology, based on the work of Tomsett et al. [249] and Onuoha [250] will be given here.

When a pellet of copper/aluminium alloy is leached by the procedure of Section 4.2.3, aluminium is selectively dissolved and the outer rim of the alloy is replaced by a uniform layer of porous copper. As the leaching proceeds, the width of the copper layer is increased until eventually the alloy particle is converted to a particle of porous copper. At incomplete stages of leaching however, the particle consists of a core of alloy surrounded by a largely uniform layer of porous copper. Pore diameters within 0 0 this copper layer are generally of the order 300 A to 400 A.

The four catalysts in Table 8.1 represent alloy particles at different stages of leaching. The thickness of copper rim can be calculated from the relationship [250]

l = Rp[t-(f)l/3] - B {8.2) 194.

where 1 = thickness of outer rim

Rp = particle radius

a = mass ratio of Al to Cu in catalyst

B = mass ratio of Al to Cu in alloy.

The composition data of Table 8.1 can now be used to calculate values of 1/Rp for each of the catalysts. These values, shown in Table 8.2 represent varying degrees of pore growth and indicate that an increase in copper surface area is associated with the growth of the porous copper layer. The relatively high BET surface areas of the partially leached catalysts can be attributed to the contribution of high surface area alumina which resides in the catalyst pores at intermediate stages of leaching [249,251].

TABLE 8.2. Evaluation of 1/Rp {Equation 8.2)

Mass Ratios (Al/Cu) Copper Surface Area Alloy {B) Catalyst

0.923 0.007 18.1 1 1.0 0.046 8.6 0.68 1.0 0.256 2.6 0.37 1.0 0.567 1.2 0.17

The residual conversion levels of the catalysts in Table 8.2 are relatively insensitive to copper surface area (Figure 8.6) and hence pore development of the catalyst indicating that, at the point where stable activity is achieved, the copper surface area in the pores makes no contribution to catalytic activity. Instead, it is proposed that the copper surface area external to the pores 195. is irrunune from deactivation and is responsible for residual activity. Once the leaching process has commenced, the surface area on the outer rim of the particles consists of pure copper and is resistant to further leaching. This fraction of the copper surface area then, should remain essentially constant, and simil­ arities in residual activity amongst Raney copper catalysts with widely differing copper surface areas, is not unexpected.

If residual activity simply reflects the availability of a small fraction of undeactivated copper surface area, then the magnitude of an activation energy calculated using a deactivated catalyst should reflect chemical reaction control. In Figure 8.5, a steady conversion level of approximately 10 percent is obtained after 50 minutes on line when a typical Raney copper catalyst is deactivated at three different temperatures. The differences in space velocity however represent a doubling in rate for each 20°C increase in temperature, or an activation energy of approximately 75 kJ mole-1• This is similar to the value of 80 kJ mol-l obtained in Section 6.4 for copper chromite catalysts over the same temperature range, and so the residual activity of the deactivated Raney copper catalyts is apparently limited by chemical reaction control.

The magnitude of the copper surface area that is external to the pores and unaffected by deactivation can readily be calculated. For the deactivation profiles of Figure 8.5, a residual conversion level of approximately 10 percent is obtained after 100 minutes on line. Substitution into Equation 4.9 yields a reaction rate of 2.3xlo-5 mole s-lg-1, which when divided by the specific activity 196.

of pure copper, established in Section 8.1 to be 2.5xlo-5 mole s -1 m-2 , yields a value of 0.9 mg2 -1 for the undeact1vated. copper surface area. Clearly, only a small fraction of the total copper surface area, as might be accounted for by the surface area exter­ nal to the pores and near the pore mouths can readily account for the observed residual activity.

In Figure 8.11, an attempt has been made to visualize the surface of the deactivated catalyst. The metal surface area within the pores is covered with polymer and is effectively inac­ tive. The fraction of the surface area external to the pores, and possibly just inside the pore rim is effectively 11 cleaned 11 by the passage of reacting gases, and is entirely responsible for residual activity. Polymer removal by depolymerization and diffusion would be in equilibrium with the production of fonnaldehyde by methanol dehydrogenation, and subsequent polymerization.

Residual activity would be critically dependent on this equilibrium which would in turn be sensitive to changes in reaction conditions. A decrease in methanol concentration would shift the equilibrium towards depolymerization resulting in an increase in active copper surface area. This is reflected in Figures 8.4 and 8.8 where a 50 percent decrease in methanol concentration leads to a marked increase in the residual conversion. An increase in space velocity would also remove more polymer counteracting the decrease in conversion that would otherwise be expected. This is al so evidem: in Fig.u.re 8.4. The effect of temperature wil 1 be com­ plicated as a higher temperature might be expected to increase the rate of depolymerization. In Figure 8.5, the increased yield

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Products Products

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I} I} FIGURE FIGURE 198. of formaldehyde expected at higher temperatures should be counte­ racted by the increase in space velocity and the earlier estimation of activation energy is probably as accurate as any that could be carried out.

The increase in residual conversion evident for higher concentrations of water in the methanol feed indicates that water may be capable of maintaining a higher active surface area than might be expected by scavenging the surface of polymerizable precursors. Moorjani et al. [86] have proposed such a contribu­ tion of water to the activity of copper catalysts for ethanol dehydrogenation (Equation 2.24).

The deactivation of Raney copper catalysts appears to be quite consistent with the idea of surface coverage by polymeriza­ tion of formaldehyde. Although the deactivation of supported copper catalysts after extended periods on line has also been discussed in terms of formaldehyde polymerization, it is not irrmediately obvious why Raney copper catalysts should deactivate at such dramatically faster rates. It might be expected for instance that the behaviour of Raney copper and silica-supported copper should be similar due to the inert nature of the silica support. In Figure 8.12, the relationship between conversion and selectivity for a Raney copper catalyst is compared with the - equivalent trends for two of the copper/silica catalysts of Figure 7.11. The data for Raney copper were obtained during a typical deactivation (Figure·B.1). Apparently, even silica has some activity for the decarbonylation of methyl formate as evidenced by the lower selectivities of the silica-supported catalysts at 199.

0 I"')

w C, 0 z N -;,-,:: a: __, :c u ><: z :c lu 0 C, z -:c 0 -U) - 0:::: a: a: UJ a: u u lu > Q.. -..J -..J z Q.. 0 C -en -en 0 u a: u 'lu ~ - >- lu' lu Q.. Q.. A.. z 0 A.. a: 0 a: u u ' • B ...

0 0 0 U) 0 U) 0 IJ> 0 - C) C) CD CD " " li.l ..l!IAI1J313S

FIGURE 8.12. The relationships bet1.;een selectivity and conversion for Raney copper and copp?.r/ silica catalysts. 220°C 200. high conversions. It might be speculated that silica also has some capacity for the breakdown of polymeric material. As the differences in selectivity are only evident at higher levels of conversion, it seems more likely that the physical nature of the catalysts is of greater importance in determining relative stability.

The growth of a high molecular weight polymer should require the presence of a large number of adsorbed formaldehyde molecules in close proximity to each other, and hence an extensive, uninterrupted copper surface area. A Raney copper catalyst should satisfy this criterion as activity is confined to the layer of porous copper. A catalyst of copper powder might be expected to behave similarly and it was noted that the copper powder catalyst of the previous chapter also exhibited sharp deactivation {Figure 7.3). In the case of a supported catalyst however, copper surface is broken up by the surface area of the support and so formalde­ hyde molecules would be consumed rapidly in the formation of methyl formate rather than in polymer growth. Furthermore, the crystallitesizesof the supported copper catalysts, shown in Table

0 7.2 to be generally in excess of 100 A, imply that much of the copper surface area is external to the small catalyst pores. Removal of polymer by the flow of reacting gases will be enhanced under these conditions.

8.4. Conclusions

Raney copper catalysts, which combine high copper surface area with the high specific activity of pure copper, should have been especially active for the dehydrogenation of methanol to 201. methyl formate. However, despite exhibiting high initial activity, the catalysts rapidly deactivated to a level of residual activity inferior to that already reported for a range of supported copper catalysts.

Deactivation is consistent with the coverage of active sur­ face by the growth of formaldehyde polymers. The catalyst pores appear to be most affected by deactivation while surface area external to the pores is largely unaffected and is responsible for residual activity. 202.

CHAPTER 9

REACTION MECHANISM

9.1. Introduction

The mechanism of methanol dehydrogenation to methyl formate was discussed in Section 2.3.2 and it was evident that there was some uncertainty as to the exact nature of the reaction.

It is generally believed that the first step in the reaction, the dehydrogenation of adsorbed methanol to formaldehyde according to

(9.1) is rate controlling.

Miyazaki and Yasumori [122] passed methanol (0.03 atmospheres) and deuterium over copper wire at 280°C. No exchange was observed suggesting that methanol was non-dissociatively adsorbed and that reaction proceeded via simultaneous removal of two hydrogen atoms. This idea supports the earlier view of Fuderer-Leutic and Brihta [47] who proposed the following general mechanism for the dehydro­ genation of alcohols on copper:-

H H I I R-C-0 ~ R-C-0 (9.2) I I I I H H H H * * 203.

H H I t

R-C-0 4 2H* + R-C=O (9.3) I I H H * *

(9.4)

A very similar mechanism has been proposed by Miyamoto and Ogino [118] to explain kinetic isotope effects in the dehydrogenation of ethanol on molten indium.

Borisov et al. [102] have suggested that the dehydrogenation of methanol on copper by reaction 9.1 might take place in two elementary steps involving the elimination of one, then another hydrogen atom. The complete mechanism might be written as

(9.5)

(9.6)

CH 30(a) 4 HCHO + H(a) (9.7)

On strong dehydrogenation catalysts such as platinum and nickel, the hydroxylic hydrogen atom of various alcohols is readily exchanged with deuterium, suggesting that dissociative adsorption does in fact take place [117,252]. Dissociative adsorption has also been proposed to explain the decomposition of methanol on zinc oxide [253] and alkaline earth metal oxides [254]. In studies of methanol oxidation on an oxygen-activated copper 110 surface, Wachs 204. and Madix [97] showed that adsorbed methoxide (CH30(a)) was the most abundant surface species at temperatures as low as 180 K. At 365 K, decomposition to adsorbed formaldehyde and hydrogen took place by the reaction

CH30(a) + HCHO(a) + H(a) (9.7)

Patterson et al. [255] have studied the exchange reaction between deuterium and various alcohols on copper. Cleavage of the aC-H and 0-H bonds were believed to take place in two distinct steps with the latter being rate controlling in alcohol dehydrogenation.

Somewhat more controversy surrounds the mechanism whereby formaldehyde is converted to methyl formate. Miyazaki and Yasumori [122] showed methyl formate to be the sole reaction product when formaldehyde was passed over copper wire at 150-180°C, and proposed methyl formate to be formed by the dimerisation reaction of two formaldehyde molecules, given by

2HCHO + HCOOCH 3 (9.8)

The reaction was believed to occur by the mechanism illus­ trated in Figure 9.1. The dimerisation mechanism was also proposed by Ai [129] to account for the formation of methyl formate in the oxidation of methanol to formaldehyde (Equation 2.29) on various binary oxide catalysts. The importance of catalyst basicity in promoting the dimerisation reaction was stressed.

Recently, Takahashi et al. [131] have studied the mechanism of formation of methyl formate from formaldehyde over copper powder 205.

H • 0 H-- t : II ____. C=O 2 H- C- H + 23" H·- C-H II 0. . lU

H » H I i I H-c~o · H-C-0 • • • • • I I H·-C-H H C-H II II 0 0 •

Figure 9.1. Dimerisation mechanism for methyl formate formation. 206. and copper/silica catalysts at 130°C using partial pressures of methanol and fonnaldehyde in the range 0.01 to 0.02 atmospheres.

The reaction of CD30H with formaldehyde yielded methyl formate with an isotopic composition indicative of HCOOCD3, and it was suggested that the dehydrogenation of methanol to methyl formate proceeded via a hemi-acetal intennediate by the mechanism illustrated in Figure 9.2, and not by formaldehyde dimerisation. The overall reac­ tion could be written as

(9.1)

CH30H + HCHO + HCOOCH3 + H2 (9.1) 2CH30H + HCOOCH3 + H2 (1.2)

In support of their claim, it was shown that ethyl formate, not methyl formate, was virtually the sole product of the reaction between formaldehyde and ethanol. It was also noted that the transesterification reaction

(9.10) did not proceed to a significant extent.

In this thesis, the dehydrogenation of methanol to methyl formate has been carried out over practical catalysts, at tempera­ tures in excess of 180°C, and using high concentrations of methanol

(> 0.3 atm). Under these conditions, the conclusions of Takahashi et al. [131] and Miyazaki and Yasumori [122] may not necessarily 207.

+ L I 0 - 0 m~ rr, ::r: I I ::I:u LJ LJ 00 0 0 -'- rr, \ I \ I 0 I LJ LJ LJ I \ I \ :::c 0 0 0 :r: :r: :C I rr, LJ '-....__/ :r: :c LJ + t - 1~ I ~ 0 :c rn I LJ :c 00 LJ \ I L LJ -0 I \ :c ~ 0 L rn 0 :r: LJ :crn Iu + oo 0 \/ I LJ LJ :r: i±

FIGURE 9.2. Hemi-acetal mechanism for methyl formate formation. 208. be valid, as these authors frequently employed very low temperatures

{130°C) and partial pressures of methanol (< 0.03 atm). In this section of the thesis, the results of experiments with formaldehyde, deuterium, unlabelled and isotopically labelled methanol will be used to gain insight into the mechanism of methanol dehydrogenation to methyl formate under practical reaction conditions.

9.2. Isotope Effects in Catalysis

Isotopic labelling has found widespread use in the study of the mechanisms of heterogeneously catalysed chemical reactions. Ozaki [256] has extensively reviewed the application of isotopic labelling to a wide range of catalytic reactions. Due to the ready availability of deuterium {2H2), most isotopic labelling involves substitution of 2H for the hydrogen atoms of organic molecules.

As well as affecting product distribution, isotopic labelling can affect the thermodynamics and kinetics of a chemical reaction due to the difference in zero-point energy between a normal organic molecule and its deuterated form. Melander [257] has shown how isotopic effects can be predicted once reaction rate and thermoydy­ namic equilibrium constants are expressed in terms of the partition functions of participating species. In the case of a kinetic isotope effect, the value of the ratio r0/rH is estimated where rH,rD represent reaction rates without substitution and with deuter­ ium substitution respectively. If the kinetic isotope effect is expressed in this manner, components of the partition function that are difficult to evaluate, such as statistical weights,symmetry numbers and moments of inertia, simply cancel out, and the ratio of rates can be expressed as 209.

where: h = Planck's constant k = Boltzmann's constant T = Temperature (K)

"H = vibrational frequency of a bond linking hydrogen to another atom (s -1 ) "o = vibrational frequency of a bond linking deuterium to another atom (s-1)

Values of "H are known and in general, "o can be calcu­ lated from the relationship [256].

"H v = - (9.12) D 2½

For instance in the breaking of C-H(C-D) or 0-H(O-D) bonds, as might occur in methanol dehydrogenation

2950 cm-l "cH = -1 "co = 2085 cm -1 "OH = 3650 cm -1 "OD = 2580 cm

Comparison of predicted and experimental kinetic isotope effects can be used in establishing a reaction mechanism. Miyamoto and Ogino [118] have studied the relative rates of dehydrogenation of isotopically labelled ethanol on indium and have proposed a 210.

reaction mechanism similar to that described by Equations 9.2, 9.3 and 9.4. Schwab and Watson [258] have used isotopically labelled formic acid to test the validity of various reaction mechanisms for the decomposition of that compound on a number of metals, including copper. Kinetic isotope effects have also been important in studies of the oxidation mechanism of olefins [259,260].

A number of factors can complicate the prediction of a kinetic isotope effect. In studies of the Fischer-Tropsch synthe­ sis over ruthenium catalysts, Kellner and Bell [261] have shown how thermodynamic isotope effects can alter the equilibrium constants of the elementary surface reactions that comprise the overall mechanism. The shift in equilibrium changes the concentration of surface species and hence the observed rate of reaction can be affected. Ozaki et al. [262] have stressed the importance of thermodynamic isotope effects in the synthesis of ammonia over iron catalysts. Recently, Davis et al. [263] have calculated the expec­ ted magnitude of kinetic and thermodynamic isotope effects for the rearrangement reactions of hydrocarbons over platinum. Experimen­ tal observations were consistent with the predicted results.

The relative strengths of adsorption of isotopic species can also be significant. Wilson [264] has speculated that kinetic isotope effects in the hydrogenation of carbon monoxide might be influenced by the preferential chemisorption of D2 over H2, as well as thermodynamic isotope effects. Soller [265] has shown o2 to be more strongly adsorbed on copper than H2, though the rate of adsorption was faster for hydrogen. 211.

9.3. Experimental

The flow reactor system and the procedures for feeding methanol, formaldehyde and the isotopes, has been described in Section 4.1.1., along with the procedure for sample collection. Sources of, and purity of the isotopes and associated reagents are listed in Table 4.3.

To avoid undue back reaction that would complicate the results of isotope experiments, it was desirable to keep conversions below 10 percent. As the syringe pumps used to feed labelled metha­ nol to the reactor could only be operated at low flowrates (< 0.1 cm3min- 1), it was generally necessary to carry out the experiments at 180°C. The use of a low reaction temperature also minimized the extent of the decarbonylation reaction of methyl formate which is given by

(9.13)

This was important as the formation of unlabelled methanol by reac­ tion 9.13 could have severely hampered analysis of the isotope experiments. At 180°C and a methanol partial pressure of 0.3 atmospheres, equilibrium conversion of methanol to methyl formate is approximately 40 percent. Experiments with formaldehyde vapour were also carried out at 180°C to minimize side reactions, most notably formaldehyde decomposition by the reaction

HCHO , CO + 2H 2 (9.14)

A copper chromite catalyst (Harshaw 0203) was chosen for most of the experiments with methanol isotopes. The catalyst had 212. a B.E.T. surface area of 16.3 mg2 -1 and a copper surface area, as determined by the nitrous oxide reaction technique of Section 4.3.2.2, of 10.8 mg2 -1 (Table 5.1). The low total surface area of the support of this catalyst (5.5. m2g-1) was particularly desirable if exchange reactions that might be support-catalysed were to be minimized. To estimate the extent of exchange due to support, the sample of cupric chromite (CuCr2o4) prepared in Section 4.2.1., was employed for one experiment. Although the B.E.T. surface area of this material (60 m2g-1) was very high as compared with the surface area of the copper chromite catalyst, it was felt that

Cucr2o4 best represented the support phase that would be encountered. To provide completely support-free copper, a Raney copper catalyst was prepared by the technique of Section 4.2.3. B.E.T. and copper surface areas of the catalyst were 18.6 and 18.1 m2g-l respectively. Copper content was determined by AAS to be 99.3 weight percent.

The Raney copper catalyst was also used for experiments with formaldehyde vapour, along with a Harshaw 1808 copper chromite catalyst with a 8.E.T. surface area of 28.4 m2g-l and a copper surface area of 6.8 m2g-l (Table 5.1).

The liquid fractions collected from each experiment involving isotopes were analysed by mass spectrometry and n.m.r. to assess the degree and nature of isotopic substitution in methanol and methyl formate molecules. Mass spectra were recorded using an AEl MS12 instrument. The vapour above a degassed sample of collec­ ted liquid was admitted and the relative amounts of d0 through d4 methyl formates were determined from parent peaks with mass to charge ratios (m/e) from 60 (d 0 ) to 64 (d4). Fragmentation was so 213. small that corrections needed were less than 1 percent. In a few cases, the isotopic composition of the methanol was also determined. However, even with minimal ionization energies, significant correc­ tions were necessary for fragmentation, and the reliability was not as high. Gas samples collected during isotope experiments were also analysed by mass spectrometry to obtain the relative amounts of H2, HO and o2. Sensitivity for these three components was low however, and there was some signal instability due to a marginal pumping speed and slight but variable exchange within the instru­ ment itself. These factors were allowed for by standardization using a standard mixture of H2 and o2 and averaging signals out over a period of time.

1H, 2H and 13c n.m.r. spectra were obtained using a Varian XL 200 Fourier transform spectrometer. Resonances due to methyl and hydroxyl (in methanol) and methyl and aldehyde (in methyl formate) hydrogen or deuterium were well separated. Corrections were needed for 13c satellites. A 20 second pulse delay was imposed to allow for relaxation and the relative areas of peaks within each spectrum were obtained by integration. Normalization between 1H and 2H peaks was achieved by including a small amount (ea. 5 percent) of a

1:1 mixture of c6H6 and c6o6 as an internal standard in each analysis. A typical set of calculations is shown in Appendix IX. The 1Hand 2H n.m.r. spectra associated with the data in Table 9.5 are shown in Figures 9.3 and 9.4. 214.

Cl

HO

u,

OHJ ---==-======::::::

CJ

Figure 9.3. 1H n.m.r. spectrum (Table 9.5). 215

00

Ln

OOJ

CJ

Figure 9.4. 2H n.m.r. spectrum (Table 9.5). 216.

9.4. Isotope Experiments

9.4.1. Results

The feasibility of carrying out experiments with methanol isotopes under practical reaction conditions was first evaluated by passing CH3oo over a Harshaw 1808 copper chromite catalyst at 200°C. Results are summarized in Table 9.1.

In agreement with the findings of Miyazaki and Yasumori

[122], the methyl formate produced was overwhelmingly HCOOCH 3. However, whereas these authors were able to obtain solely HD, analysis of the gas phase concentration in Table 9.1 shows a mixture of H2, o2 and HD, and a value of D/(H+D) significantly lower than the value of 0.5 that would be expected. This latter result is most likely due to dehydrogenation of CH 30H, and as shown in Table 9.1, a high concentration of CH30H (13.7 percent) is evident in the liquid as a result of exchange at the hydroxyl group of methanol.

It will be shown later (Table 9.4) that the dehydrogenation of CH30H proceeds twice as rapidly as the dehydrogenation of CH30D. The expected yields of H2 and HD can thus be estimated from the concen­ trations of these two species via

Yield H ~~~2 = 2x13.7 = • 34 Yield D2 79.6 T which corresponds to a value of D/(H+D) of 37.5 percent. This value agrees closely with that in Table 9.1.

The concentrations of H2, D2 and HD can be termed "statis­ tical" or equivalent to the thermodynamically expected values. 217.

TABLE 9.1. Results of CH30D dehydrogenation

Catalyst - 1808 Copper Chromite 0.5 g Temperature - 200°c LHSV of Methanol - 11 h-l Mole Fraction of Methanol in Feed - 0.7 Conversion - 1.7 percent Liquid Analysis (ex n.m.r.) 1. Methanol before reaction

CH30D 99.8 percent CH30H 0.18 CH2DOD 0.02 2. Methanol after reaction

CH30D 79.6 percent CH30H 13.7 CH2DOD 5.7 CH2DOH 1.0 3. Methyl formate HCOOCH3 82.6 percent DCOOCH3 8.4 HCOOCH2D 8.2 DCOOCH2D + HCOOCHD2 2.5 Gas Analysis (ex mass spec.) expected (statistical) D2 10 percent 12.2 percent HD 49 45.5

H2 41 42.3

D/(H+D) = 0.35 218.

(For evaluation see Appendix X). The disagreement between this result and the findings of earlier experiments [122] is undoubtedly related to the large differences in reactant concentrations between the two sets of experiments. As mentioned in Section 9.1, Miyazaki and Yasumori [122] used methanol partial pressures of 0.03 atmospheres and observed first order kinetics. Under these condi­ tions, adsorbed D and H atoms would be comparatively widely spaced and the two atoms liberated according to

(9.15) would be more likely to combine with each other than with Hand D atoms from other molecules. Under the conditions used to obtain the data for Table 9.1, zero order kinetics are obeyed (Section 7.5) and adsorbed methanol molecules are tightly packed together so the Hand D atoms of neighbouring methanol molecules are more likely to combine.

To further investigate the nature of H2;o2 equilibration, a mixture of methanol, H2 and D2 was passed over the same copper chromite catalyst. The results of analysis are surrmarized in

Table 9.2. A 11 statistical 11 mixture of D2, H2 and HO was once more obtained indicating that equilibration can also proceed via dissoci­ ative adsorption of gaseous molecules. As the catalyst surface would be saturated with adsorbed methanol, few sites would be available for adsorption of H2, D2 or HO and therefore the equili­ bration process must be extremely rapid. 219.

TABLE 9.2. Results of CH 30H/H2/D2 exchange

Catalyst - 1808 Copper Chromite 0.5 g Temperature - 210°c LHSV of Methanol - 24 h-l

Mole Fractions1 - YCH 0H = 0.61; YH = 0.20; v0 = 0.19 3 2 2 Conversion - 3.7 percent

Liquid Analysis 1. Methanol (ex n.m. r.)

CH30H 88.9 percent CH30D 9.9 CH 2DOH 1.2 2. Methyl formate (ex mass spec.)

HCOOCH3 91.5 percent DCOOCH3 + HCOOCH2D 8.5 Gas Analysis (ex mass spec.) expected (statistical) 6.1 percent 5.3 percent 33.3 35.4 60.6 59.3

3 -1 1. Total gas flowrate (210°C) = 407 cm min . 220.

Tables 9.1 and 9.2 show that considerable exchange has taken place into both methanol and methyl formate, with the hydroxyl group of methanol apparently being particularly vulnerable. In an attempt to minimize exchange, and hence clarify the results of analysis procedures, subsequent experiments were performed at a lower temperature (180°C), and using a Harshaw 0203 copper chromite catalyst. As mentioned in the previous section, this catalyst had a particularly low total surface area, and a high ratio of copper surface area to total surface area, important factors if exchange was support-catalysed.

Table 9.3 summarizes the results obtained when a mixture of

CH 30H and o2 was passed over the 0203 copper chromite at 180°C. Considerable exchange into the methanol molecule is still evident and is occurring predominantly at the hydroxyl position, in agree­ ment with the results of Tables 9.1 and 9.2. To estimate the activity of the support for exchange, an equivalent mixture of

CH 30H and o2 was passed over cupric chromite (Cucr2o4 ) at 200°C. Exchange into the methyl position was negligible(< 0.01 percent) and as no dehydrogenation activity was evident, it was suspected that this form of exchange might arise via the interaction of 0 with adsorbed formaldehyde. A typical sequence of reactions could be

(9.1)

HCHO + HO + CH200H (9.16) 221.

Cupric chromite did have considerable activity for exchange into

the hydroxyl group of methanol (D/(H+D) = 5.4 percent). However, this result is clouded by the high surface area of the material (60 m2g- 1) and the possible existence of surface protons arising from the acid leaching process used in the preparation procedure (Section 4.2.1).

To finally clarify the activity of copper for exchange, a mixture of CH30H and D2 was passed over a Raney copper catalyst that had been allowed to deactivate to a steady level of conversion. As Table 9.3 shows, the extent of exchange over Raney copper is in excess of that obtained over the copper chromite catalyst, presum­ ably due to the lower space velocity employed in the former case. Clearly copper is very active for hydroxyl exchange and the extent of substitution at the hydroxyl position is high as compared with that at the methyl position of the methanol molecules. As will be discussed in Section 9.5, the existence of a significant concentra­ tion of adsorbed methoxide (CH30(a)) is indicated.

In view of the findings of Chapter 8, it might be expected that exchange could arise from interaction of D with the high molecular weight surface species that was proposed to give rise to deactivation. A more complex pattern of substitution than was obtained would surely have been expected however. Furthermore, as the pattern of exchange of Raney copper mirrors that obtained over the copper chromite catalysts, it is very reasonable to expect that the exchange activity of copper is being observed. Equilibration of H2, o2 and HD occurs over Raney copper but not completely over the copper chromite catalyst indicating that the deactivated 222.

TABLE 9.3. Results of CH 30HJD2 exchange at 180°C

Catalyst Copper Chromite Raney Copper 0203 (1 g) (2 g}

LHSV of Methanol (h-1) 26 11

Mole fractions1 CH30H 0.84 0.82 D2 0.16 0.18 Conversion (percent} 1.6 1.4

Liquid Analysis (ex n.m.r.} (D/(H+D)) Unreacted methanol (percent)

"methyl 11 0.26 1.9 "hydroxyl" 7.1 19.8

Gas Analysis2 (percent) (ex mass spec.) H2 14.7 (7.9) 34.2 (33.6) HD 26.7 (40.4) 46.8 (48.7)

D2 58.6 (51.7) 19.0 (17.7) 0/(H+D) 72.0 42.1

1. Total gas flows: 611 cm3min-l (0203); 540 cm3min-l (Raney). 2. "Statistical" values in parentheses. 223. surface might not be totally impervious to dissociative adsorption of H2, D2 or HD. The lower space velocity employed over Raney copper could also be responsible for this result.

As the influence of support on the exchange reaction was secondary to that of copper, the change in copper chromite catalyst from Tables 9.2 to 9.3 did little to influence the extent of substitution. Nevertheless, subsequent experiments were carried out using the 0203 copper chromite. The results of the kinetic isotope experiments are presented in Table 9.4. Zero order kine­ tics are expected on the basis of earlier results (Section 7.5) and reaction rates were evaluated by the procedure outlined in

Section 4.1.4. The value r0/rH is the ratio of reaction rates for isotopically labelled and unlabelled methanol, and the minimum value of 0.12 obtained for co3oo corresponds to an eightfold decrease in rate due to deuterium substitution. This result was quite unprecedented as no kinetic isotope effect of similar magni­ tude could be found in the literature. Also shown in Table 9.4 are the expected values of r0/rH obtained via Equation 9.11 for the following possibilities of reaction rate control.

1. 0-H bond breaking; 2. C-H bond breaking; 3. Simultaneous C-H, 0-H bond breaking.

The implications of the observed and predicted kinetic isotope effects will be discussed in Section 9.4.2.

The experiments represented in Table 9.4 required the use of a nitrogen carrier gas to ensure an even flow of methanol at the 224.

TABLE 9.4. Kinetic isotope experiments at 180°C1

2 rD/rH Isotope LHSV Conversion Experimental Predicted (h-1) (%) 1 2 3 CH30H 9.6 9.5 1.00 1.00 1.00 1.00 CH300 4.8 9.8 0.52 0.18 1.00 0.18 CD30H 4.8 4.7 0.25 1.00 0.25 0.25 co3oo 4.8 2.2 0.12 0.18 0.25 0.05

1. Catalyst - Harshaw 0203 copper chromite (0.5 g)

2. YAlcohol = 0.35 (in nitrogen) low flowrates employed. Chromatographic separation of nitrogen and carbon monoxide could not be achieved and so it was not possible to assess the extent of methyl fonnate decarbonylation by reaction

9.13. However, the ratio of CO/N2 was determined by mass spectro­ metry to be less than 1/200 for the dehydrogenation of CH3oo in Table 9.4, so the generation of unlabelled methanol by reaction 9.13 was assumed to be insignificant.

Two isotope experiments were specifically carried out to gain insight into the mechanism whereby methyl formate is obtained from formaldehyde. In each case, the 0203 copper chromite was used and a temperature of 180°C employed. Conversions were relatively high to ensure an adequate concentration of methyl formate for accurate analysis by n.m.r. In the first experiment, the extent of transesterification by reaction 9.10 was assessed by passing a

3/1 mixture of co30H and HCOOCH3 over the catalyst. Table 9.5 sulTITlarizes the input and output isotope distributions of methanol and methyl formate. Clearly, the transesterification reaction proceeds rapidly over the catalyst at 180°C. 225.

TABLE 9.5. Results of 11 CD30H/HC00CH/ experiment

Catalyst - 0203 copper chromite 1 g Temperature - 180°C 1 Mole fractions - YCD30H = 0.23; YHCOOCH3 = 0.078 Conversion - 15 percent Liguid Analysis (ex n.m.r.) (D/(H+D)) 1. Methanol before reaction Percent D in "hydroxyl" = 0 percent

Percent D in "methyl 11 = 98.2 2. Methyl formate before reaction

Percent D in 11 aldehyde 11 = 0 percent

Percent D in 11 methyl 11 = 0 percent 3. Methanol after reaction

Percent D in "hydroxyl 11 = 19.6 percent

Percent D in "methyl 11 = 78.5 percent 4. Methyl formate after reaction

Percent D in 11 aldehyde 11 = 17.1 percent

Percent D in "methyl 11 = 79.9 percent

Gas Analysis

H2 74.6 percent HD 24.8 02 0.6 D/(H+D) = 13 percent

3 . -1 1. Balance N2; Total gas flow= 149 cm m1n 226.

The transesterification should be relatively unaffected by deuterium substitution in the methyl positions of methanol and methyl formate, as these groups do not participate in the reaction, and an equilibrium constant approximately equal to unity would be expected. It can readily be shown that the values of D/(H+D) in Table 9.5 for substitution into the methyl groups of methyl formate and methanol correspond closely to the values of 0.75 expected if the transesterification reaction was to reach equilibrium. Analysis by mass spectrometry confirmed that the formation of co3 accounted for over 90 percent of deuterium substitution in the methyl posi­ tions of both methanol and methyl formate, and that the quantities of CD3, CH2D and CHD2 were similar in each species. These results indicate that the transesterification reaction proceeds at a much more rapid rate than methanol dehydrogenation. The low values of D/(H+D) obtained in Table 9.5 for the gas analysis can also be related to transesterification. As Table 9.4 shows, dehydrogenation of

CH30H formed in reaction 9.10 can be expected to yield H2 at four times the rate that CD30H would yield HD, hence contributing to an apparent excess of H. Exchange of deuterium into the hydroxyl position of methanol (19.6% in Table 9.5) will also decrease the value of D/(H+D).

In a related experiment, a 3/1 mixture of co3oo and CH30H was passed over the catalyst. The transesterification and dehydro­ genation of each species would be expected to influence the product distribution and as can be seen in Table 9.6, at least five iso­ topic species of methyl formate are present. Values of D/(H+O) for substitution into the methyl positions of methanol and methyl 227.

TABLE 9.6. Results of "CD300/CH30H" experiment

Catalyst - 0203 copper chromite 1 g Temperature - 180°C Mole fractions1 - YCH 0H = 0.083; Yeo OD= 0.25 3 3 Conversion - 10.7 percent

Liquid Analysis (ex n.m.r.) (D/(H+D)) 1. Methanol after reaction Percent D in "hydroxyl" = 72.5 percent Percent D in "methyl" = 73.6 percent 2. Methyl formate after reaction Percent Din "aldehyde" = 35.3 percent

Percent Din "methyl 11 = 80.5 percent

Ex mass spec. d = 13.1 percent 0 dl = 8.3 d2 = 5.2 d3 = 50.2 d4 = 23.2

Gas Analisis H2 = 37.7 percent HO = 39.7

D2 = 22.6 D/(H+D) = 42.5

1. Balance N2; Total gas flow= 152 cm3min-l 228.

formate approximate the value of 0.75 expected for equilibration of the transesterification reaction.

The results of this and the previous experiment (Table 9.5) are relevant to the formation of methyl formate from formaldehyde, and so will be discussed, along with the results of experiments with formaldehyde, in Section 9.5.2.

9.4.2. Discussion

Comparison of experimental and predicted values of r0/rH in Table 9.4 indicates that the observed kinetic isotope effects can­ not be explained by simply assuming rate control due to one elementary reaction step. If rate is assumed to be controlled by

0-H bond breaking, then the kinetic isotope effect due to co3 substitution cannot be explained. Although the kinetic isotope effect observed for CD30H can be accurately predicted if C-H bond breaking is assumed to be rate controlling, the effect of 0-0 substitution must be accounted for. The reaction mechanism outlined in Equations 9.2, 9.3 and 9.4, involving simultaneous cleavage of C-H and 0-H bonds, also fails to predict the experimental trends observed in Table 9.4.

It was suspected initially that the results of Table 9.4 might be complicated by a change in the thermodynamic equilibrium constant of the reaction, due to isotopic substitution. Under the conditions used throughout this thesis, methanol dehydrogenation is an equilibrium reaction and the maximum theoretical conversion of methanol rarely exceeds 40 percent (Figure 6.1). As the equili­ brium constant is necessarily related to the ratio of forward and reverse rate constants, a decrease in the equilibrium constant will 229.

have a corresponding effect on reaction rate. Cant et al. [266] have estimated thermodynamic equilibrium constants for the dehydro­

genation of CH 30H, CH30D, CD30H and co3oo, from the partition functions of the participating species. Isotopic substitution was found to increase the magnitude of the equilibrium constant by up to a factor of four times, though the influence on the equilibrium yield of methyl formate was not great. Clearly, the decreases in rate evident in Table 9.4 cannot be attributed to an isotope effect of this type.

It is possible that a thermodynamic isotope effect could affect the reaction rate by changing the concentration of a parti­ cular surface species. As was mentioned in Section 9.2, this type of isotope effect has been demonstrated for Fischer-Tropsch synthesis [261], the synthesis of ammonia [262], and for the hydro­ genolysis of hydrocarbons over platinum [263]. In the case of methanol dehydrogenation, the rate controlling step might involve the dehydrogenation of adsorbed methanol to formaldehyde according to

(9.1)

the dehydrogenation of adsorbed methanol to adsorbed methoxide according to

(9.6) or the dehydrogenation of adsorbed methoxide by

CH 30(a) ~ HCHO + H(a) (9. 7) 230.

It is possible then that the kinetic isotope effects might be

related to the concentrations of CH 30H(a), CH30(a), HCHO(a) and H{a). As hydrogen is only adsorbed very weakly on copper, rate control due to the desorption of hydrogen must inmediately be dis­ missed. Since the dehydrogenation reaction has been shown to obey zero order kinetics {Section 7.5) the catalytic surface must be considered to be saturated with adsorbed methanol. Under these conditions, it is hard to envisage that the observed kinetic isotope effects could be related to differences in concentration of adsorbed methanol amongst the isotopes. If adsorbed formaldehyde existed in a high concentration, then the extent of exchange into the methyl position of methanol by the reaction

HCHO(a) + HD + CDH 20H (9.16) would be expected to be greater than is observed in Figure 9.3. A much higher rate of exchange into the hydroxyl position of methanol is evident however, and the existence of a significant concentration of adsorbed methoxide is indicated.

As was mentioned in Section 9.1, CH30(a) is known to be the dominant surface species when methanol is adsorbed onto an oxygen­ activated copper 110 surface. Sexton [267] has used high resolution electron energy loss spectroscopy {EELS) to study the surface species present on a clean copper 100 surface. Methoxide ions were readily formed below room temperature, and decomposed to gaseous products at 370 K. 231.

The concentration of CH30(a) might be expected to influence the rate of either of reactions 9.6 and 9.7. Since Table 9.3 shows the rate of hydroxyl exchange to be far in excess of the rate of methanol dehydrogenation, it would appear that reaction 9.6 is highly unlikely to be rate controlling. It is more likely that this reaction is at equilibrium and that reaction 9.7 is rate controlling. Under these conditions, it is feasible that isotopic substitution could affect reaction rate by altering the equilibrium of reaction

9.6, and hence the small but critical concentration of CH30(a). To evaluate the effect of isotopic substitution, two equilibria must initially be considered

Kl CH30H{a) ! CH30(a) + H(a) (9.6)

(9.17)

If reactions 9.6 and 9.17 are interconnected, then the two equili­ bria must be solved simultaneously. However, this situation is unlikely. The results in Table 9.2 have already been shown to imply that the equilibration of reaction 9.17 proceeds much more rapidly than the rate of dehydrogenation. Furthermore, the results of

CH30HJD2 exchange over Raney copper (Table 9.3) show equilibration of o2, HD and H2, yet hydroxyl exchange is obviously incomplete. It appears reasonable then to consider the equilibrium of reaction 9.6 in isolation. 232.

Under these conditions, the concentration of adsorbed methoxide can be expressed according to

(9.18}

given that [H(a}] = [CH30(a}] To estimate the effect of deuterium substitution at the hydroxyl position, it is necessary to evaluate the ratio

[CH30(a)] ex CH OD R=--- 3 (9.19) [CH30(a)] ex CH OH 3 where K10 , K1H are the values of K1 (Equation 9.6) applicable for dehydrogenation of CH30D and CH30H respectively. It should be noted that, for the assumption of zero order kinetics, the terms [CH30H(a)] and [CH300(a}] that would otherwise feature in Equation 9.19, cancel out. To evaluate K10/KlH it is necessary to express K10 and K1H in terms of the partition functions involved in the equilibria of Equation 9.6 [256] i.e. K1D QD QCH30H KlH = (QH} . (QCH30D}exp - [(ZH-ZD) - (ZCH30D-ZCH30H}]/RT (9.20)

where QN = partition function of species N ZN = zero point energy of species N R = gas constant T = temperature 233.

(Note that the partition functions for the two methoxide species are identical and cancel out).

Cant et al. [266] have estimated the values of K10JK1H from the expected translational, rotational, vibrational and zero point energies of adsorbed H, D, CH30H and CH30D. Maximum and minimum values of 0.79 and 0.31 were obtained and corresponded to values of between 0.89 and 0.56 for R in Equation 9.19. Allowing for some inaccuracy in these estimations, it does appear that the substitu­ tion of deuterium in the hydroxyl position will certainly lower reaction rate by decreasing the concentration of adsorbed methoxide in reaction 9.7.

It is proposed then that the results of Table 9.4 represent a combination of kinetic and thermodynamic isotope effects. The rate controlling step appears to be the dehydrogenation of surface methoxide by reaction 9.7, as the value of r0/rH for co30H can be accurately predicted from Equation 9.11 for the case where C-H bond breaking is rate controlling. The substitution of deuterium at the hydroxyl position gives rise to a thermodynamic isotope effect that acts to decrease the concentration of adsorbed methoxide in the rate controlling step. As Table 9.4 shows, the magnitude of the thermo­ dynamic isotope effect (for CH30D) is approximately half that of the kinetic isotope effect (for co30H). As would be expected, the lowest rate of dehydrogenation is obtained for co3oo, which combines both types of isotopic hinderance. 234.

9.5. Experiments with Formaldehyde

9.5.1. Results

To evaluate the extent to which copper catalysed the dimeri­ sation of formaldehyde to methyl formate by reaction 9.8, formaldehyde vapour {30 percent in helium) was passed over a copper chromite catalyst {Harshaw 1808) and a Raney copper catalyst at 180°C. Results are presented in graphical form in Figures 9.5 and 9.6.

Steady conversion levels were obtained after 15 minutes on line, and as Figures 9.5 and 9.6 show, conversion of formaldehyde to methyl formate over both catalysts was very high. This reflects earlier conclusions which point to the dehydrogenation of methanol to formaldehyde as the rate controlling step {Sections 6.4 and 9.4.2). Differences in carbon monoxide levels between the two catalysts can be accounted for by the activity of the copper chro­ mite catalysts for methyl formate decarbonylation by the reaction (Sections 6.4 and 7.5)

{9.13)

In the absence of catalyst, bulk reaction of formaldehyde was less than 2 percent.

Clearly, under the conditions of high formaldehyde concen­ tration, the dimerisation reaction is predominant. The approximate equivalence of carbon monoxide and methanol levels in Figures 9.5 and 9.6 suggests that methanol formed by reaction 9.13 does not further participate in the synthesis of methyl formate. 235.

1

• 8

• 6 ~ HCCJCJCH3

[!) CH30H z: 1D...... A,. CCJ t--- . 4 u er: X HCHCJ 0:::: lL w _J ~ :::::E: . 2

0

0 5 10 15 20 TIME(MINUTES)

FIGURE 9.5. Product distribution for HCHO/He over 1 copper chromite 1808. 180°C; 300 cm3min- ; YHCHO = 0.3; 1 g catalyst 236,

1

• 8

. 6 ~ HCCICICH3

[!] CH3CIH z D,._. .t. CCI I- . 4 u CI 0:: LL w _J D :::E: . 2

0

0 5 10 15 20 TIME (MINUTES)

FIGURE 9.6. Product distribution for HCHO/He over Raney copper. 180°C; 300 cm3min-1; YHCHO = 0.3; 1 g catalyst 237.

It was noted that, after exposure to formaldehyde vapour at 180°C for 20 to 30 minutes, the activity of the two catalysts for methanol dehydrogenation to methyl formate (Equation 1.2) was no more than 50 percent of that obtained with fresh catalysts. This result reinforces earlier conclusions that the deactivation of catalysts for methanol dehydrogenation was related to polymerization of formaldehyde.

9.5.2. Discussion

The results of the previous section would tend to indicate that the dimerisation reaction of formaldehyde according to

2HCHO + HCOOCH3 (9.8) was responsible for the production of methyl formate during methanol dehydrogenation. The conclusions of Ai [129], who has emphasized the importance of acidic and basic sites in the promotion of formal­ dehyde dimerisation, are questioned as neither Raney copper nor copper chromite could be expected to have strong acidic or basic functions.

It must be noted however, that the experimental conditions used to obtain Figure 9.5 and 9.6 may yield a catalytic surface unrepresentative of that found in methanol dehydrogenation. Under the conditions of high formaldehyde partial pressure, adsorbed formaldehyde molecules would be expected to be in close proximity to each other, and the probability of dimerisation by reaction 9.8 would be high. In the case of methanol dehydrogenation, the catalytic surface is saturated with adsorbed methanol (and methoxide) 238.

and the ratio of methanol sites to formaldehyde sites would undoub­ tedly be very large. Under these conditions, the hemi-acetal mechanism depicted in Figure 9.2 and represented by the reaction

{9.9)

could reasonably be expected to occur.

As mentioned in Section 9.1, Takahashi etal. [131] observed an increase in conversion of formaldehyde to methyl formate over copper, with the addition of methanol over the temperature range 130-180°C. However the possibility of methanol dehydrogenation according to

{1.2) could not be discounted. Attempts to reproduce the experiments of these authors in the temperature range of interest {180-220°C) were unsuccessful due to the extremely high rate of formaldehyde conver­ sion by reaction 9.8.

Takahashi et al. [131] also observed HCOOCD3 to be the predominant reaction product when a mixture of CD30H and HCHO was passed over copper. This result could only be explained by reaction 9.9 as the dimerisation reaction {9.8) would be expected to yield

HCOOCH3 or a mixture of HCOOCH3, HCOOCD2H, DCOOCH2D and DCOOCD3 from the reactions of HCHO and DCOO. At the low partial pressures

{Pco OH' PHCHO < 0.02 atm) and temperatures {130°C) used, the 3 transesterification reaction 239.

(9.10) did not proceed to a significant extent.

The experiment involving dehydrogenation of co3oo and CH30H (Table 9.6) was aimed at resolving the mechanism of methyl formate formation under practical reaction conditions. If the dimerisation reaction was predominant, then a mixture of HCOOCH3 (d 0 ), DCOOCH2D (d2}, HCOOCD2H (d2} and DCOOco3 (d4) would be formed. No d1 or d3 species would be·expected. If, on the other hand, methyl formate was produced by a hemi-acetal mechanism, a mixture of HCOOCH 3 (d0 ), DCOOCH3 (d1), HCOOCD3 (d3) and DCOOCD3 (d4) might be obtained. No d2 methyl formate would be expected. The absence of either d2 or d1 and d3 methyl formate would have theoretically made it possible to distinguish between the mechanisms represented in Figures 9.1 and

9.2. In the case of the dimerisation mechanisms however-, d1 and d3 can be formed from the transesterification reactions

(9.10) and (9.21)

The breakdown of the hemi-acetal intermediate in Figure 9.2 can yield CHD20H and CH2DOH which can transesterify to d2 methyl formate according to

(9.22)

(9.23) 240.

Obviously, identification of the correct reaction mechanism will be dramatically complicated in the presence of transesterification.

The results outlined in Tables 9.5 and 9.6 have already been shown to indicate that the transesterification reaction proceeds to equilibrium. This finding is in conflict with the results of Takahashi et a 1. [ 131] but can be accounted for by the considerable difference in reaction conditions between the two sets of experi­ ments. At the low partial pressures of reactants employed by those authors, transesterification would be hindered by the distance between adsorbed molecules. However, under the conditions used to obtain the results of Table 9.5, adsorbed molecules would be closely packed and transesterification more likely to occur. The lower temperature (130°C) used by Takahashi et al. might also restrict transesterification.

Although d0 , d1, d2, d3 and d4 methyl formate were all represented in Table 9.6, it was hoped that the d1 and d3 methyl formate could be attributed to transesterification, and that the level of d2 methyl formate, though small, might be important. How­ ever, the production of d2 methyl formate by exchange, as evident in Table 9.1, makes it impossible to establish the significance of this result.

In su111T1ary, it would appear that the mechanism whereby formaldehyde is converted to methyl formate cannot be identified under practical reaction conditions due to transesterification. The results of Takahasi et al. [131] must be accepted at this stage, and the mechanism of Figure 9.2, involving hemi-acetal formation, is assumed to be valid. 241.

9.6. Conclusions

The mechanism of methanol dehydrogenation to methyl formate has been investigated using isotopically labelled methanol, deuterium and formaldehyde.

Dehydrogenation to formaldehyde is believed to take place according to

(9.5}

(9.6)

CH30(a) + HCHO + H(a) (9.7)

2H{a) + H2 {9.24)

The kinetic isotope effect observed for CD30H indicates that reaction 9.7 is the rate controlling step, though deuterium substitution at the hydroxyl position of methanol also influences rate. The latter effect is most likely due to a thermodynamic isotope effect which serves to decrease the concentration of adsorbed methoxide in the rate controlling step.

The formation of methyl formate from formaldehyde which can be accounted for by either of the reactions

2HCHO -+ HCOOCH3 (9.8)

(9.9) 242.

is very rapid and does not influence the overall rate of reaction. Unfortunately, transesterification makes distinction between the two mechanisms impossible, though the mechanism described by Figure 9.2 appears to be the most feasible. 243.

CHAPTER 10

CONCLUSIONS AND RECOMMENDATIONS

In this investigation various copper catalysts have been shown to be active and selective for the dehydrogenation of methanol to methyl formate at atmospheric pressure, and over the temperature range 180-240°C.

Three commercial copper chromite catalysts were characteri­ zed by elemental analysis, nitrogen adsorption to determine total surface area, reaction with nitrous oxide to determine copper surface area, X-ray diffraction and thermal gravimetric analysis. Despite the high copper loadings and low total surface areas of the catalysts, relatively high metal surface areas and small copper crystallite sizes were obtained after reduction in pure hydrogen at 220°C. The reduction process is described by the reactions

(10.1)

{10.2)

{10.3)

Reaction 10.1 occurs very rapidly, accompanied by the formation of cuprous chromite by reaction 10.2. Copper chromite catalysts prepared via precipitation of copper ammonium chromate undergo further, though incomplete, reduction by reaction 10.3 to yield a catalytic surface of copper supported on chromia. However, in the 244. case of a catalyst prepared by calcination of copper and chromium oxides, cuprous chromite is not reduced, presumably due to a shield­ ing effect of the large excess of copper.

These findings confirm the results of earlier studies where X-ray photoelectron spectroscopy (XPS) had been used to monitor catalyst reduction during high pressure, liquid phase hydrogenation. It is reconmended that XPS be utilized in any further research into the nature of copper chromite catalysts.

Preliminary thermodynamic calculations indicated that the suitability of copper chromite catalysts for dehydrogenation should be evaluated at atmospheric pressure and with the use of a pure methanol feed.

The presence of chromia in the copper chromite catalysts appears to reduce both activity and selectivity, with the latter effect being attributed to the activity of chromia for the decar­ bonylation of methyl formate. Catalyst deactivation, which is pronounced in the absence of chromia, is attributed to the polymer­ ization of formaldehyde. Chromia is believed to prolong active life by promoting breakdown of the carbonaceous species.

Under the conditions established for catalyst comparison, the dehydrogenation reaction obeys zero order kinetics. In the absence of support interaction, activity is proportional to copper surface area, in spite of very large differences in copper disper­ sion amongst the various catalysts. The comparatively low specific activities obtained for the copper chromite catalysts, when compared with those for catalysts of pure copper or copper 245. supported on silica, cannot be attributed to mass transfer limita­ tions or overestimation of metal surface area. Instead a detrimental effect of chromia on the activity of copper, as confirmed by the particularly low specific activity of copper supported on chromia, is proposed, and attributed to a form of electronic interaction between copper and copper-chromium oxides. The selectivity of the various copper catalysts is strongly related to the activity of the support for the decarbonylation of methyl formate.

Despite possessing low specific activity, the copper chromite catalysts emerge as being most effective for the dehydrogenation reaction due to their high bulk density and relatively high selec­ tivity, though it is reconunended that the stability of the catalysts over extended periods of time be evaluated. Any kinetic modelling that might be envisaged should be carried out using one of the more stable copper chromite catalysts. It might be of interest, using the standard conditions established in this thesis, to investigate the effectiveness of the various mixed oxide catalysts that have been proposed in Patent literature to be suitable for the dehydro­ genation reaction.

Given the high specific activity of pure copper, Raney copper catalysts, which possess particularly high metal surface areas, should have proven very effective. Various Raney copper catalysts were tested, and though initial activity was often very high, all underwent dramatic deactivation to a level of residual activity that was largely independent of copper surface area, and inferior to the activity of many of the supported catalysts previously tested. To explain this phenomenom, it is proposed that 246.

the interior surface of the catalyst pores is deactivated by the growth of formaldehyde polymer. Residual activity may be attributed to the fraction of the copper surface area external to the pores, that should remain essentially constant irrespective of the width of the porous copper rim around the leached particles. The slow, but complete, catalyst reactivation that was observed suggests control of reactivation rate by a depolymerization process.

Techniques for studying the surface of a deactivating Raney copper catalyst under conditions of practical interest are very limited. It is suggested that the deactivation of a series of Raney copper catalysts of very small particle size be examined. As the surface area external to the pores of such catalysts would be expec­ ted to constitute a greater fraction of total copper surface area, high levels of residual activity might be obtained.

The mechanism of the dehydrogenation reaction was investi­ gated by means of experiments with deuterium, isotopically labelled methanol, and fonnaldehyde. These experiments indicate that the initial dehydrogenation occurs according to the reaction sequence

CH30H + CH30H(a) (10.4)

CH30H{a) + CH30(a) + H(a) (10.5)

CH30(a) + HCHO + H{a) (10.6)

Reaction 10.6 is shown to be rate controlling and influenced by a large kinetic isotope effect due to the substitution of deuterium 247. atoms for methyl hydrogen atoms in the methanol molecules. A thermodynamic isotope effect due to deuterium substitution at the hydroxyl position of methanol, that acts to decrease the concentra­ tion of adsorbed methoxide in the rate controlling step, has also been identified. As both kinetic and thenoodynamic isotopic hinderance acts to decrease reaction rate, the overall effect of isotopic substitution is very large.

Identification of the mechanism whereby formaldehyde is converted to methyl formate is not possible due to a transesterifi­ cation reaction that equilibrated rapidly and randomized the effect of deuterium substitution. The weight of evidence, however, favours the formation of methyl formate via a reaction mechanism involving hemi-acetal formation. The feasibility of further experiments with isotopes is dependent on preparation of a copper catalyst that will minimize the extent of transesterification.

Although the synthesis of methyl formate by the copper­ catalysed dehydrogenation of methanol has considerable industrial viability, it is strongly recommended that the techniques and find­ ings of this thesis be extended to the study of formaldehyde synthesis by methanol dehydrogenation. Though temperatures in excess of 550°C would be required, the dehydrogenation process has the potential of producing formaldehyde at a much higher selectivity than is currently possible with the conventional oxidation process. The central problem of such a research project would appear to be the preparation of a copper catalyst possessing relatively stable activity under high temperature operation. 248.

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243. Niiyama, H., Tamai, S., Kim, J., and Echigoya, E., J. Japan Petrol. Inst., 24(5), 322 (1981). 244. Dowden, D.A., Catal. Rev.,~' 1 (1972). 245. SocieteI I Normande de Produits Chimiques, Brit. Pat. 629312, Sept. 16 (1949) 246. Yuekl'son, I.I., Terekhim, P.M., and Kvashnina, S.P., Khim. Prom. (Moscow), 43(2), 83 (1970); C.A. 73:14090y. 247. Innes, W.B., in Errmett, P.H. ed., "Catalysis", Reinhold, N.Y., 1961, Vol. 1, p.299. 248. Sherwood, T.K., Pigford, R.L., and Wilke, C.R., "Mass Transfer", McGraw Hill, Tokyo, 1975, p.19. 249. Tomsett, A.D., Young, D.J., and Wainwright, M.S., J. Electrochem. Soc., In Press. 250. Onuoha, N.I., Ph.D. dissertation, University of New South Wales, 1982. 251. Robertson, S.D., and Anderson, R.B., J. Catal., 23, 286 (1971). 252. Bond, G.C., "Catalysis by Metals", Academic Press, London, 1962, p.220. 253. Ueno, A., Onishi,, T., and Tamaru, K., Trans. Faraday Soc., 67, 3585 (1971). 254. Okamoto, Y., Imanaka, T., and Teranishi, S., Bull. Chem. Soc. Japan, 45, 3207 (1972). 255. Patterson, W.R., Roth, J.A., and Burswell, R.L., J. Amer. Chem. Soc., 93(4), 839 (1971). 256. Ozaki, A., "Isotope Effects in Heterogeneous Catalysis", Academic Press, N.Y., 1977. 257. Melander, L., "Isotope Effects on Reaction Rates", Ronald Press, N.Y., 1960. 264.

258. Schwab, G.M., and Watson, A.M., Trans. Faraday Soc., 60, 1833 {1964). 259. Adams, C.R., and Jennings, T.J., J. Catal., l, 549 (1964). 260. Cant, N.W., and Hall, W.K., J. Catal., 52, 81 (1978). 261. Kellner, C.S., and Bell, A.T., J. Catal., 67, 175 (1981). 262. Ozaki, A., Taylor, H., and Boudart, M., Proc. Roy. Soc., 258A, 47 (1960). 263. Davis, S.M., Gillespie, W.D., and Somorjai, G.A., J. Catal., 83, 131 (1983). 264. Wilson, T.P., J. Catal., 60, 167 (1979). 265. Soller, T., Goldwasser, S., and Beebe, R., J. Amer. Chem. Soc., 58, 1703 {1936). 266. Cant, N.W., Tonner, S.P., Trimm, D.L., and Wainwright, M.S., Submitted to J. Catal., April, 1984. 267. Sexton, B.A., Surf. Sci., 88, 299 (1979). 265.

APPENDICES 266.

APPENDIX I

PUBLICATIONS REGARDING THE BASE-CATALYSED CARBONYLATION OF ALCOHOLS The following articles have been removed from the digital copy of this thesis. Please see the print copy of the thesis for a complete manuscript.

Title: THE BASE-CATALYSED CARBONYLATION OF HIGHER ALCOHOLS Authors: S. P. TONNER, D. L. TRIMM, M. S. WAINWRIGHT Journal: Journal of Molecular Catalysis, 18 (1983) 215 · 222

Title: Solubility of Carbon Monoxide in Alcohols Authors: Stephen P. Tonner, Mark S. Wainwright, and David L. Trimm Journal: JOURNAL OF CHEMICAL & ENGINEERING DATA 1983 Vol. 28, no.1, p. 59-61 278.

APPENDIX II

THERMODYNAMIC YIELD OF METHYL FORMATE BY FORMALDEHYDE CONDENSATION - CALCULATIONS

2HCHO -+ HCOOCH 3 (1)

- -1 tiH298 = 118 kJ mole

tiS298 = -137 J mol-lK-l . tiGT i' tiH298 - TtiS298 (2) and tiGT = -RT ln kp (3)

.. . Combining (1) and (2)

ln kp = 14i93 - 16.5 = ln ky (at 101 kPa) (3)

From Equation (1), total moles (NT) at conversion = X:-

2HCHO -+ HCOOCH 3

N = X = 2-X (4) T 2-2X Mole Fractions given by:-

= 2-2X = X (5,6) YHCHO 2-X YHCOOCH3 2-X

2 . 2X-X (7) .. ky = 2-2X

Substitute ky from (3) and solve (7) for X by trial and error. 279.

APPENDIX I II

BASIC Computer Program for calculation of conversion, selectivity, mole fractions, rate, LHSV and pressure. 280. 10 INPUT• CATALYST •;et 20 INP~t~ ~ASS OF CATALYST 9 ";N 30 INPUT" TEMPERATURE C •;TT -40 T=TT+273 50 INPUT" FLOURATE "1/"in ";L 60 INPUT" PRESSURE at" ";P 70 INPUT" co ";co 80 INPUT" C02 ";C02 90 INPUT" H20 ";H20 100 INPUT" HEOH ";NEOH 110 INPUT" HEF ";HEF 120 CO=C0/.8 130 C02=C02/. 87 140 HEF=HEF/1.2 150 THEF=CO+C02+2*HEF 160 TC=HEOH+2•NEF 170 TH=CO+C02+H20+NEOH+NEF 180 Xl=TNEF/TC 190 X2=2•HEF/TC 200 SEL=X2/X1 210 F2=(1-Xt•.5)/(1+X1*.5) 220 YCO=F2•CO/TH 230 YC02=F2•C02/TH 240 YH20=F2•H20/TH 250 YHEOH=F2•HEOH/TN 260 YNEF=F2•NEF/TN 270 LPRINT" CATALYST "Ct 280 LPRINT 290 LPRINT" NASS 9 "N 300 LPRINT" TEMPERATURE C "TT 310 LPRINT" FLOURATE "1/"in "L 320 PP=P-•.5 330 LPRINT" PRESSURE at" "PP 340 LPRINT 350 LPRINT 360 F=L•.0004167 370 R=F*Xl/N 380 LHSV=L•.8•60/H 390 LPRINT" NOLE FRACTIONS" 400 LPRINT 410 LPRINT" HEOH"TAB(20)"CO"TAB(34l"NEF"TAB<48l"H20"TAB(62l"C02• 420 LPRINT 430 LPRINT TAB(4l YHEOH TAB(17) YCO TAB(31l YNEF TAB(45) YH20 TAB(59) YC02 440 LPRINT 450 LPRIIH 460 LPRINT" CONVERSION % "Xl 470 LPRINT 480 LPRINT" SELECTIVITY "SEL 490 LPRIIH 500 LPRINT" LHSV ( /HRl "LHSV 510 LPRINT 520 LPRINT" RATE Nole/s/9 "R 530 LPRINT 540 Lf'RINT !550 Lf"fil N r 560 INPUT" CHANGE PRESSURE & FLOURATE (YES=1,NO=O>•;H !:i70 IF Ne:: 1 THEN 50 !:i80 GOTO /0 '.:i90 STOP 281.

APPENDIX IV

Preparation of Copper/Silica Catalyst by 11 incipient wetting 11 technique - Sample Calculation.

Required Catalyst - 10 g of 5 weight% Cu/Si02

Water required to 11 wet 11 support - 12 ml for 10 g of support.

Concentration of Stock Solution - 166.7 g Cu/1000 ml.

Required mass of copper in catalyst -

10 X 5 = 0.5 g. 100

Volume of Stock Solution to be made up to 12 ml -

= 0.5 g x 1000 ml = 3 ml of Stock Solution 166.7 g 282.

APPENDIX V

CALCULATION OF COMPOSITION OF REDUCED COPPER CHROMITE CATALYSTS

Chemical Reactions

CuO -+ Cu+ 0 (1)

Cu+ Cucr2o4 -+ cu2cr2o4 (2)

cu 2cr2o4 -+ 2Cu + Cr2o3 + 0 (3)

2BaCr04 -+ Cr2o3 + 2Ba0 + 3x0 (4)

Molecular Weights (g)

Cu = 63.5 (Ml) CuO = 79.5 (M2) 231.5 (M3) cucr2o4 = 295 (M4) Cufr2o4 = Cr2o3 = 152 (MS) BaO = 153 (M6) BaCr04 = 253 (M7) 0 = 16 (MB)

Harshaw 1808 (Basis 100 g)

Composition before reduction

Total topper = 35.6 g (a)

= 23.9 g = 19.1 g Cu (b)

= 60.3 g = 16.5 g Cu (c)

Inerts = 15.8 g (I) 283.

Composition after reduction

Weight Loss (W) = 5.67 g (Table 5.7)

Ml Ml Cu= Wx - - Cu Cr O x - (5) M8 2 2 4 M4

(Figure 5.5)

Cu = 22.5 g M3 CuCr2o4 = [a-Cu] x M (6) 1

= 47.8 g

M M Cr203 = [Cu-b] X 2 + [BaO] X 2 (7) M1 2 M6

Bao= o

Percentage on I+ Cu+ Cucr2o4 + cu2cr2o4 + cr2o3 + BaO for Table 5.8.

Girdler G-22 (Basis= 100 g)

Composition before reduction

Total Copper = 33.7g (a) CuO = 25.9 g = 20.7 g Cu (b) Cucr2o4 = 47.4 g = 13.0 g Cu (c) BaCr04 = 18.5 g = 11.1 g BaO after reduction Inerts = 8.2 g 284.

Composition after reduction

Weight loss = 8.99 g Weight loss due to Equation 4 = 1.75 g

True weight loss = 7.24 g

Cu (from Equation 5) = 28.7 g

Cu 2cr2o4 = 0 (Figure 5.5)

Cucr2o4 (from Equation 6) = 18.2 g Bao = 11.1 g

Cr2o3 (from Equation 7) = 24.7 g Percentage as previously.

Harshaw 0203 (Basis= 100 g)

Composition before reduction

Total copper = 59.4 g (a)

CuO = 64.5 g = 51.5 g Cu (b) = 28.9 g = 7.9 g Cu (c)

Inerts = 6.6 g

Composition after reduction

Moles cu2cr2o4 = Moles Cucr2o4 (before reduction)

Cu 2cr2o4 = 36.8 g

Weight loss = 13 g

Cu (from Equation 5) = 43.7 g 285.

Cucr2o4 = 0 Bao = o

Cr2o3 = 0

Percentage as previously. 286.

APPENDIX VI

CALCULATION OF EQUILIBRIUM YIELDS OF METHYL FORMATE

Inerts + 2CH30H + HCOOCH3 + 2H2 + Inerts

Total moles

(NT) = (M0 -M0 x) + (F0+M0 X/2) + (H0+M0x) + I

= Mo+ Fo + H0 + M0 X/2

where:- X = Conversion of methanol Mo = Initial moles of methanol Fo = Initial moles of methyl formate Ho = Initial moles of hydrogen I = Moles of inerts Ky = (Y 2y H2 HCOOCH3)/YCH OH2 3 where = (M0 -M0 X)/NT YcH30H

YHCOOCH3 = (F0+M0 X/2)/NT

YH = (H0+M0X)/NT 2 YI = I/NT From thermodynamic data:­ -4790 ln Ky = -T- + 6.87xln T-0.0055 x T - 32.70

BASIC computer program solves for X by Interval Halving. 287.

10 DEF FNXCP,T)•EXPC-4790/l+6.87•LOG(T)-.0055•T-32.7)/P 20 DEF FNXCX)•FNK(P,T>•2•NT•(H-N•X>·2-(2•F+N•X>~CH+N•X>~2 30 DEF FNNCX>•I+N+F+H+H•X/2 ~O INPUT•PRESSURE ATH";P 50 INPUT"TEHPERATURE C ";T 60 TaT+27J 70 INPUT"HOLES HETHANOL";H 80 INPUT"HOLES FORHATE";F 90 INPUT"HOLES H2";H 100 INPUT"HOLES INERT";I 110 XR:a1 120 Xl=O 130 NT•FNN(XL> 140 YL=FNX CXL> 150 NT=FNNCXR> 160 YR=FNXCXR> 170 X=CXR+XL>/2 175 Y=FNX(X) 180 NT=FNNCX> 190 IF YL+Y=O GOTO 220 200 IF YL•Y H TAB<20> F TAB(34> H TAB(48) I 400 LPRINT 410 LPRINT 420 LPRINT" EQUILIBRIUH CONVERSION= "X 430 LPRINT 440 LPRINT 450 LPRINT" EQUJLIBRIUN HOLE FRACTIONS" 460 LPRi NT 470 LPRINT" HEOH"TAB(20>"HEF"TAB(34l"H2•tABl48)"lNERT" 480 LPRINT 490 LPRINT TAD(4) YH TAB(20) YF TAB(34) YH TABC48) YI 500 LPRINT 510 LPRINT 520 LPR lNT 530 LPl

APPENDIX VI I

SAMPLE CALCULATIC~ OF DISPERSION CATALYST - COPPER CHROMITE 1808

0 = surface atoms of copper total atoms of copper

Copper Surface Area (Table 5.1) = 6.8 m2g-l

Copper after reduction (Table 5.8) = 0.24 g (g catalyst)-l

6.8 m2 x 1.46 x 1019 atoms g m2 D = 23 .24 _g__ x 6x10 atoms g cat. 63.5 g

= 0.044 289.

APPENDIX VI II

EXPERIMENTAL DATA FOR EVALUATION OF COPPER CATALYSTS 290.

·catalyst - Copper Chromite 1808 Mole Fractions LHSV Pressure Conversion Selectivity (atm) (%) (%) (h-1) co CH30H HCOOCH 3

20.2 1.05 .028 .70 .075 21.5 82.6 50.9 1.10 .009 .83 .046 11.2 88.7 74.9 1.15 ·.006 .87 .036 8.4 90.1 110.9 1.20 .003 .91 .026 6.1 90.6 150.8 1.25 .002 .93 .021 4.8 90.5 190.7 1.30 .'002 .94 .015 3.5 89.2

. :catalyst - Copper Chromite 0203

Mole.Fractions LHSV Pressure Conversion Selectivity (atm) (%) (%) (h-1) co CH 30H HCOOCH 3

14.9 1.03 .032 .54 .121 36.0 86.6 26.4 1.04 .013 .65 .102 26.3 91.6 52.8 1.07 .004 .79 .066 15.1 94.9 81.6 1.11 .002 .85 .046 10.1 95.6 115.2 1.16 .001 .89 .034 7.4 95.6 153.6 1.25 .001 .92 .025 5.4 94.7 196.8 1.30 .001 .94 .019 4.1 94.4 291.

:catalyst - Copper Chromite G-22

Mole.Fractions LHSV Pressure Conversion Selectivity (atm) (%) (%) (h-1) co CH30H HCOOCH 3

14.9 1.03 .046 .63 .086 27.2 78.7 26.4 1.03 .021 • 71 ~074 20.6 83.9 52.8 1.06 ~007 .82 .050 12.2 89.9 81.6 1.10 .004 .87 .037 8.5 91.2 115.2 1.15 .003 .90 .028 6.4 90.6 153.6 1.22 .002 .93 .018 4.1 90.1 196.8 1.30 .002 .94 · .016 3.7 90.5

Catalyst - Copper/Chromia

Mole Fractions LHSV Pressure Conversion Selectivity (~tm) (%) (%} (h-1) co CH30H HCOOCH 3

14.9 1.02 .038 .87 .016 7.9 45.9 26.6 1.05 .011 .94 .009 3.4 52.8 53.1 1.08 '.005 .97 .005 1.8 59.7

.... -- -- 292 .

. Catalyst - Copper/Magnesia

Mole Fractions LHSV Pressure Conversion Selectivity (atm) (%) (%) (h-1) co CH30H HCOOCH 3

28.9 1.2 .060 .62 .069 26.1 69.7 49.8 1.05 .026 .80 .055 14.9 81.0 89.6 1.07 .010 .85 .046 10.7 90.6 181.8 1.15 .003 .91 .024 5.5 93.8

.Catalyst - Copper/Silica (High)

Mole.Fractions LHSV Pressure Conversion Selectivity (atm) (%) (%) (h-1) co CH30H HCOOCH3

29.7 1.04 .033 .62 .100 28.2 86.2 --- 56.l 1.04 .015 .72 .082 20.3 91.7 105.8 1.10 '.006 .83 .051 11.8 92.0 161. .1 1.20 .004 .88 .035 8.0 92.6 218.9 1.30 .003 .91 .026 5.8 92.4 293. :catalyst - Copper/Silica {Low)

LHSV Pressure · · Mole.Fractions Conversion Selectivity (atm) (%) (%) (h-1) co CH30H ~COOCH3

26.4 1.05 .002 .85 .046 10.2 95.8 52.8 1.11 .001 .93 .023 4.9 96.1

80.7 1.23 ' - .95 .014 3.1 94.8 14.9 1.04 .006 .74 .079 18.6 94.7

.. :catalys - Copper.Silica (Medium)

LHSV Pressure Mole.Fractions Conversion Selectivity (atm) (%) (%) (h-1) .co CH30H HCOOCH 3

26.4 1.04 .005 .73 .081 19.0 95.4 52.8 1.10 .002 .85 .046 10.1 96.5 81.6 1.19 ·.001 .91 .029 6.2 95.7 109.5 '1.30 .001 .94 .019 4.2 95.7

·catalyst - Copper Silica (Ion fxchange)

LHSV Pressure Mole.Fractions Conversion Selectivity (atm) (%) (%) (h-1) co CH 30H HCOOCH3

24.1 1.08 .052 .58 .094 31.1 78.2 52.9 1.07 .018 . 71 .080 20.5. 89.0 105.6 1.15 ·.oos .82 .053 12.1 94.1 158.1 1.25 .002 .87 .039 8.8 94.1 294.

APPENDIX IX

Calculation of the degree of deuterium substitution (D/(H+D)) for Table 9.6. Areas (from Integration) 1 2 Compound lH H/C6H6 2H H/C6D6 Benzene 48.48 6.04 Methyl formate ( a ) 11 a 1dehyde 11 5.53 .114 0.475 .079 (b) "methyl 11 8.04 .166 3.87 .641 Methanol (c) "hydroxyl 11 43.18 .891 13.254 2.194 (d) "methyl" 111.44 2.299 36.23 5.998

R (molar ratio c6H6Jc6o6) = 0.934

2 H/C6o6 D/(H+D) =

For methyl formate:

D/(H+D)(a) = ---·7-9--­ = .426 (.79 + .934 X .114)

0/(H+D)(b) = (.641 + :;~~ x .166) = .805

For methanol

D/(H+D)(c) = 2.194 +2:~~: x .891 = .725

5.999 D/(H+D)(d) = 5.999 + .934 x 2.299 = • 736 295.

APPENDIX X

Estimation of 11 statistical 11 H2,o2 and HD.

From Table 9.1:-

H2 = 41%, o2 = 10%, HD = 49% This is equivalent to

41 + 49 = 2 66% "2

10 + 49 = 2 Now, K for

H2 + o2 t 2HD (Total moles= 2) is approximately equal to 4 (252) Therefore:-

_ (2 X 66X) 2 4 - {66-66X){34-66X) (X = conversion)

X = 0.34 (by trial and error) So:-

"2 = 66-66(0.34) = 43.6% 02 = 34-66(0.34) = 11.6% HO = 2x66 (0.34) = 44.8%