kinetic requirements for catalytic deoxygenation of ... · reaction mechanisms and kinetics,...

Kinetic Requirements for Catalytic Deoxygenation of Alkanals on Solid Brønsted Acid Sites

by

Fan Lin

A thesis submitted in conformity with the requirements for the degree of Doctor of Philosophy

Department of Chemical Engineering & Applied Chemistry University of Toronto

© Copyright by Fan Lin 2017

ii

Kinetic Requirements for Catalytic Deoxygenation of Alkanals on

Solid Brønsted Acid Sites

Fan Lin

Doctor of Philosophy

Department of Chemical Engineering & Applied Chemistry

University of Toronto

2017

Abstract

Alkanals undergo catalytic deoxygenation on solid Brønsted acid catalysts (e.g. acidic zeolites

and polyoxometalate clusters), producing light alkenes, alkadienes, or heavy aromatic species,

under atmospheric pressure and moderate temperature (473-673 K). The details about the

reaction mechanisms and kinetics, however, remain ambiguous. In this thesis, C3-C6 n-alkanals

were used as model reactants in combined kinetic measurement, chemical titration, and infrared

spectroscopic experiments, to investigate the mechanisms and kinetic requirements for alkanal

deoxygenation on solid Brønsted acid sites (H+). The reaction network for alkanal (CnH2nO)

deoxygenation consists of three kinetically coupled primary reaction pathways: (i) intermolecular

C=C bond formation, a bimolecular pathway lengthening the carbon chain via aldol

condensation and dehydration and forming larger alkenals (e.g. C2nH4n-2O and C3nH6n-4O), (ii)

intramolecular C=C bond formation, an unimolecular pathway evolving alkene (CnH2n) via

transfer hydrogenation and dehydration while preserving the carbon backbone, and (iii)

isomerization-dehydration, another unimolecular pathway directly ejecting a water molecule

from the alkanal and producing alkadiene (CnH2n-2). The secondary cyclization-dehydration

following pathway (i) produces aromatic species which provide the H atoms required for the

alkanal transfer hydrogenation during pathway (ii); pathway (i) is also kinetically coupled with

iii

pathway (iii) as both pathways share a bimolecular surface intermediate and the accessibility of

acid-base site pairs dictates the selectivities between these two pathways. In pathway (ii), the

kinetic relevant alkanal transfer hydrogenation step proceeds via hydride ion transfer from a

hydrocarbon (e.g. aromatic products) as the H-donor to the protonated alkanal as the H-acceptor,

upon forming a bimolecular carbocation transition state. The hydride transfer is favored on H+

sites under confinement of molecular dimension because the confined structure solvates and

stabilizes the transition state and lowers the activation barrier. The rate of the transfer

hydrogenation is dictated by the hydride ion affinity difference between the H-donor-acceptor

pair. The chain length of the alkanal reactant determines the stability of the enol tautomer and the

hydride ion affinity of the protonated alkanal, as a result, dictating the rates and selectivities of

these deoxygenation pathways. This mechanistic knowledge on the multiple catalytic cycles and

their kinetic and thermodynamic requirements provides the framework for predicting the rates

for larger oxygenates and hydrocarbons during alkanal deoxygenation and could guide the design

of catalyst structures to enable tuning of rates and product distributions.

iv

Acknowledgments

The author would like to express sincere gratitude to the following individuals, groups and

organizations.

Supervisor: Professor Dr. Ya-Huei (Cathy) Chin

Co-workers: Members of the Multidisciplinary Laboratory for Innovative Catalytic Science,

University of Toronto

Collaborators: Members of Technische Chemie II (Professor Johannes Lercher Group),

Technische Universität München, Germany

Supervisory Committee: Professor Dr. Charles Mims and Professor Dr. Bradley Saville

Graduate Unit: Department of Chemical Engineering and Applied Chemistry, University of

Toronto

Funding Agencies and Scholarships: Natural Sciences and Engineering Research Council of

Canada (NSERC), Valmet, Abellon CleanEnergy, Canada Foundation for Innovation (CFI);

Hatch Graduate Scholarship for Sustainable Energy Research, Ontario Graduate Scholarship, and

DAAD (German Academic Exchange Service) Scholarship for Academic Exchange

v

Table of Contents

Acknowledgments .......................................................................................................................... iv

Table of Contents ............................................................................................................................ v

List of Tables .................................................................................................................................. x

List of Figures ............................................................................................................................... xii

List of Schemes ........................................................................................................................... xxii

Preface ......................................................................................................................................... xxv

Chapter 1 Introduction to Alkanal Deoxygenation on Solid Brønsted Acid Catalysts ........... 1

References .................................................................................................................................. 7

Chapter 2 Mechanism of Intra- and Intermolecular C=C Bond Formation of Propanal on

Brønsted Acid Sites Contained within MFI Zeolites ................................................................. 9

2.1. Introduction ....................................................................................................................... 10

2.2. Experimental ..................................................................................................................... 12

2.2.1. Catalyst preparation ................................................................................................ 12

2.2.2 Catalytic rates and selectivities of propanal and 1-propanol reactions on MFI

zeolites .................................................................................................................. 12

2.2.3. Chemical titration of Brønsted acid sites ................................................................ 14

2.2.4. Temperature programmed desorption of surface intermediates after propanal

reactions on MFI zeolites ...................................................................................... 14

2.3. Results and discussion....................................................................................................... 15

2.3.1. Reaction network and product distributions during catalytic deoxygenation of

propanal on H-MFI zeolites .................................................................................. 15

2.3.2. Accessibilities of Brønsted acid site to propanal reactant and effects of acid site

density on propanal conversion rates .................................................................... 20

2.3.3. Kinetic dependencies, elementary steps, and site requirements for

intermolecular C=C bond formation of propanal on H-MFI zeolites ................... 23

2.3.4. Kinetic dependencies, elementary steps, and site requirements for

intramolecular C=C bond formation in propanal on H-MFI zeolites ................... 28

vi

2.3.5. Reversibility of the inter- and intramolecular C=C bond formation in propanal

on H-MFI zeolites ................................................................................................. 31

2.3.6. Kinetic relevance of hydrogen transfer and requirements of hydrogen for

intramolecular C=C bond formation in propanal .................................................. 34

2.3.7. Regression of rate data with the derived rate expressions for inter- and

intramolecular C=C bond formation ..................................................................... 37

2.4. Conclusion......................................................................................................................... 41

2.5. References ......................................................................................................................... 42

2.6. Appendix ........................................................................................................................... 46

2.6.1. Mass balance during propanal reaction on H-MFI ................................................. 46

2.6.2. Time on stream evolution of propanal conversion ................................................. 47

2.6.3. Determination of kinter,eff and k−inter,eff in Equation 2.9 ............................................. 48

2.6.4. Determination of kinetic parameters in Equations 2.10 and 2.11 by non-linear

regression fitting of these equations with rate data for C3H6O-C6H12 reactions

on H-MFI zeolites in Figure 2.7 ........................................................................... 49

Chapter 3 Alkanal Transfer Hydrogenation Catalyzed by Solid Brønsted Acid Sites .......... 52

3.1. Introduction ....................................................................................................................... 53

3.2. Experimental ..................................................................................................................... 55

3.2.1. Catalyst preparation ................................................................................................ 55

3.2.2 Rate and selectivity assessments .............................................................................. 55

3.3. Results and discussion....................................................................................................... 56

3.3.1. Alkanal deoxygenation pathways and the kinetic couplings of intramolecular

C=C bond formation in alkanals and dehydrogenation of aromatic products at

Brønsted acid sites ................................................................................................ 56

3.3.2. Mechanism of transfer hydrogenation between tetralins or cyclohexadienes and

protonated alkanals at Brønsted acid sites ............................................................ 66

3.3.3. Catalytic effects of alkanal molecular size and local acid site confinements on

transfer hydrogenation reactions ........................................................................... 75

3.4. Conclusion......................................................................................................................... 83

3.5. References ......................................................................................................................... 84

vii

3.6. Appendix ........................................................................................................................... 88

3.6.1. Rate equation for intramolecular C=C bond formation .......................................... 88

3.6.2. Intermolecular C=C bond formation and Tishchenko esterification of butanal

on NaH-MFI zeolites ............................................................................................ 89

3.6.3. Estimation of hydride ion affinities for protonated alkanals and carbenium ions

of H-donors ........................................................................................................... 90

3.6.4. Effects of time-on-stream on the amount of remaining H+ sites, the rate of

butanal conversion, and the reaction selectivities on H-MFI, H-FAU, and

H4SiW12O40 catalysts ............................................................................................ 93

3.6.5. Parity plots for the kinetic data ............................................................................... 95

3.6.6. Characterizations of the Brønsted and Lewis acid sites ........................................ 101

Chapter 4 Kinetic Requirements of Solid Brønsted Acid Catalyzed Transfer

Hydrogenations of Aldehyde ................................................................................................. 103

4.1. Introduction ..................................................................................................................... 103

4.2. Experimental ................................................................................................................... 105

4.2.1. Catalyst preparation .............................................................................................. 105

4.2.2 Rate and selectivity assessments ............................................................................ 105

4.2.3. Infrared spectroscopic study ................................................................................. 106

4.3. Results and discussion..................................................................................................... 107

4.3.1. Kinetic and infrared spectroscopic studies on aldehyde transfer hydrogenation

by hydrocarbons on Brønsted acid sites .............................................................. 107

4.3.2. Effects of carbon chain length on the aldehyde transfer hydrogenation ............... 117

4.4. Conclusion....................................................................................................................... 120

4.5. References ....................................................................................................................... 120

4.6. Appendix ......................................................................................................................... 122

4.6.1. Infrared spectra of H-FAU upon pyridine and butanal adsorption ....................... 122

4.6.2. Carbon distribution in the aromatic product fraction during butanal reaction on

H-FAU zeolite ..................................................................................................... 123

4.6.3. Determination of aldehyde transfer hydrogenation rate by co-feed H-donors ..... 124

viii

4.6.4. Kinetic relevance of hydrogen transfer step in aldehyde transfer hydrogenation

on Brønsted acid sites ......................................................................................... 125

4.6.5. H+ site coverage by butanal and its derivatives in infrared spectroscopic study .. 125

4.6.6. Estimation of the H+ site coverage by carbonyl group on H-FAU ....................... 127

4.6.7. Comparison of rates for intermolecular C=C bond formation with and without

H-donor incorporation ........................................................................................ 131

4.6.8. Estimation of hydride ion affinities for protonated aldehydes and carbenium

ions of H-donors ................................................................................................. 131

Chapter 5 Catalytic Pathways and Kinetic Requirements for Alkanal Deoxygenation on

Solid Tungstosilicic Acid Clusters ......................................................................................... 134

5.1. Introduction ..................................................................................................................... 135

5.2. Experimental ................................................................................................................... 137

5.2.1. Preparation and characterizations of H4SiW12O40 clusters dispersed on SiO2

support ................................................................................................................. 137

5.2.2 Rate and selectivity assessments for alkanal deoxygenation on H4SiW12O40

polyoxometalate clusters ..................................................................................... 139

5.3. Results and discussion..................................................................................................... 140

5.3.1. Catalytic pathways of alkanal deoxygenation on H4SiW12O40 tungstosilicic acid

dispersed on high surface area silica substrates .................................................. 140

5.3.2. Kinetic coupling of alkanal chain growth, deoxygenation, and isomerization-

dehydration cycles .............................................................................................. 149

5.3.3. Mechanisms for the formation of alkenals, alkenes, and dienes via primary

alkanal reactions on H4SiW12O40 clusters ........................................................... 152

5.3.4. Catalytic sequences for secondary cyclization-dehydration that form aromatics

and cycloalkadienes on H4SiW12O40 clusters ...................................................... 162

5.3.5. Effects of alkanal chain length on its deoxygenation rates and selectivities on

H4SiW12O40 clusters ............................................................................................ 167

5.4. Conclusion....................................................................................................................... 171

5.5. References ....................................................................................................................... 173

5.6. Appendix ......................................................................................................................... 177

5.6.1. Characterizations of Brønsted and Lewis acid sites on H4SiW12O40 .................... 177

ix

5.6.2. Stability of polyoxometalate at high temperature ................................................. 178

5.6.3. Coke formation during butanal reactions on H4SiW12O40 .................................... 179

5.6.4. Effects of extra-framework alumina on the butanal reaction on H-MFI zeolite ... 181

5.6.5. Temperature programmed desorption (TPD) of pyridine on H4SiW12O40 ........... 183

5.6.6. Reactions of 2,4-heptadienal on H4SiW12O40 ....................................................... 183

Chapter 6 Summary of Alkanal Deoxygenation on Solid Brønsted Acid Sites and

Perspective for Bio-oil Upgrading ......................................................................................... 185

6.1. Summary of catalytic pathways for alkanal deoxygenation ........................................... 185

6.1.1 Intermolecular C=C bond formation ...................................................................... 185

6.1.2. Intramolecular C=C bond formation ..................................................................... 186

6.1.3. Direct alkanal dehydration .................................................................................... 188

6.2. Perspective on the study of catalytic deoxygenation of bio-oil on solid Brønsted acid

catalysts ........................................................................................................................... 189

6.3. References ....................................................................................................................... 190

Appendix: Error Assessment ...................................................................................................... 192

A1. Error assessment for rate measurements and chemical titrations under identical

conditions ........................................................................................................................ 192

A2. Error assessment for rate assessments at varying alkanal pressures ............................... 195

A3. Error assessment for the measured rates of alkanal transfer hydrogenation ................... 196

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List of Tables

Table 2.1. Pyridine uptakes on H-MFI zeolites from pyridine titration carried out after steady-

state reactions in C3H6O-H2O mixturesa ....................................................................................... 31

Table 2.2. Rate parameters derived from non-linear regression fittings of rate data to rate

equations [Eqn. 2.8 (in Sec. 2.3.4), Eqn. 2.9 (in Sec. 2.3.5), Eqns. S2.2 and S2.3 (in Appendix

Sec. 2.6.4)] .................................................................................................................................... 38

Table 3.1. Rates and selectivities for butanal deoxygenation and butanol dehydration on H-MFI,

H-FAU, or H4SiW12O40 at 573 K .................................................................................................. 60

Table 3.2. The extent of promotion, ,tetralinj , ,tetralin-adj

, or ,chdj for the various reactions j

(j=Inter, Intra, Dehy, or Tish) with tetralin, tetralin-adamantane, or cyclohexadiene incorporation

during butanal deoxygenation, and the rate constant for cyclohexadiene-to-butanal transfer

hydrogenation, 4 8,C H OTH -chd

k , on H-FAU, H-MFI, and H4SiW12O40 at 573 K ............................. 69

Table 3.3. Rates for tetralin dehydrogenation and butanal hydrogenation on H-FAU and

H4SiW12O40 at 573 K .................................................................................................................... 71

Table 3.4. Transfer hydrogenation rates of butanal (, 4 8TH C H O-tetralinr ), butadiene (

, 4 6TH C H -tetralinr ), and

butene (, 4 8TH C H -tetralinr ) by tetralin on H-FAU at 573 K .................................................................. 73

Table S2.1. Parameter values derived from non-linear regression of rate data in Figure 2.7 (H-

MFI, Si/Al=11.5, 7.5 ks, 1.1×10−3

mol C3H6O·(mol H+

i·s)−1

, 1.1 kPa C3H6O in He, 473K) with

Equations S2.4 and S2.5 ............................................................................................................... 51

Table S3.1. Hydride ion affinities (HIA) for protonated alkanals (CnH2nOH+; n=3-6) and the

carbenium ions of the H-donors (R’H+) ........................................................................................ 93

Table S3.2. The amounts of Brønsted and Lewis acid sites on H-MFI, and H-FAU, and

H4SiW12O40 catalysts .................................................................................................................. 102

xi

Table S4.1. Initial total H+ site coverage by butanal and its derivatives ( Total ) and H

+ site

coverage by butanal (4 8C H O,0 ) on H-FAU before feeding H-donors. ......................................... 127

Table S4.2. Rate ratios for the pathway of intermolecular C=C bond formation in C4H8O-RDH2

feed mixture (D4 8 2,C H O R HInter -r ) to that in C4H8O feed (

4 8,C H OInterr ) on H-FAU at 573 K .............. 131

Table S4.3. Hydride ion affinities (HIA) for protonated aldehydes (CnH2nOH+; n=3-6) and the

carbenium ions of the H-donors (RDH+). .................................................................................... 133

Table S5.1. The amounts of Brønsted and Lewis acid sites on H4SiW12O40 catalysts ............... 177

Table S5.2. Amount of Brønsted and Lewis acid sites on H-MFI zeolites (Si/Al=40) and the

ammonium hexafluorosilicate treated H-MFI zeolites [H-MFI(AHFS)] (measured by FTIR study

of pyridine adsorption at 448 K) ................................................................................................. 182

xii

List of Figures

Figure 2.1. Temperature dependence of propanal conversion rates (○) and the rates of olefin

(C2=-C6

=, ×), C6H10O (2-methyl-2-pentenal, ●), C9H14O (2,3,4,5-tetramethyl-2-cyclopentenal

and isomers, ▲), aromatic (C6-C12, ■) formation during propanal (C3H6O) reactions on H-MFI

zeolites [Si/Al=11.5, 7.5 ks, 1.1×10−3

mol C3H6O·(mol H+

i·s)−1

(subscript i denotes the initial

acid site density), 1.1 kPa C3H6O in He]. ..................................................................................... 16

Figure 2.2. Carbon distributions in the effluent stream of propanal (C3H6O) reactions on H-MFI

zeolite (Si/Al=11.5) at 473 K (a), 523 K (b), 548 K (c), and 673 K (d) (7.5 ks, 1.1×10−3

mol

C3H6O·(mol H+

i·s)−1

, 1.1 kPa C3H6O in He, overallr is the overall C3H6O conversion rate). ......... 17

Figure 2.3. Desorption rate of carbonaceous species from H-MFI catalyst as a function of

temperature. The temperature programmed desorption was performed after exposure of the

catalyst to propanal (C3H6O) reactions for 960 s at 473 K (300 mg H-MFI, Si/Al=11.5, 0.0167

K∙s−1, propanal reaction conditions: 1.1 kPa C3H6O in He, 1.23×10

−6 mol C3H6O·(gcat.·s)

−1). .... 21

Figure 2.4. Overall rates (per mass of catalyst, ●) and turnover rates (per H+

i, subscript i denotes

the initial acid site density, ○) for intramolecular C=C bond formation in propanal (C3H6O) on

H-MFI plotted as a function of H+

i and Na+ concentration (473 K, Si/Al=11.5, 7.5 ks, 1.23×10

−6

mol C3H6O·(gcat.·s)−1

, 1.1 kPa C3H6O in He). .............................................................................. 23

Figure 2.5. Turnover rates for intermolecular C=C bond formation (rinter, ■) and intramolecular

C=C bond formation (rintra, ●) in propanal (C3H6O) that evolve 2-methyl-2-pentenal (C6H10O)

and propylene, respective, and the rate ratio for inter- over intramolecular C=C bond formation

(rinter/rintra, ○), plotted as a function of C3H6O pressure on H-MFI at 473 K [Si/Al=11.5, 7.5 ks,

1.1×10−3

-4.4×10−3

(mol C3H6O·(mol H+

i· s)−1

)] ........................................................................... 24

Figure 2.6. Effects of water on turnover rates for intramolecular C=C bond formation (intra

r , ●)

and intermolecular C=C bond formation (inter

r , ▲) in propanal (C3H6O) during C3H6O reactions

on H-MFI at 473 K (Si/Al=11.5, 7.5 ks, 1.1×10−3

mol C3H6O·(mol H+

i·s)−1

, 1.1 kPa C3H6O in

He). ................................................................................................................................................ 33

xiii

Figure 2.7. Effects of 3-methyl-1-pentene (C6H12) pressure on intramolecular C=C bond

formation (intra

r , ●) and intermolecular C=C bond formation (inter

r , ▲) in propanal (C3H6O)

and the rate ratio for intra- over intermolecular C=C bond formation (inter

r /intra

r , ○) during

C3H6O reactions on H-MFI catalysts at 473 K (Si/Al=11.5, 7.5 ks, 1.1×10−3

mol C3H6O· (mol

H+

i·s)−1

, 1.1 kPa C3H6O in He). .................................................................................................... 35

Figure 3.1. Rates for intramolecular C=C bond formation (Pathway 2, 4 8,C H OIntrar ) as a function of

aromatic pressure ( AromaticsP ) during butanal reactions on H-MFI [■, Si/Al=11.5, space velocity

0.0033-0.013 mol butanal (mol H+ s)

–1], H-FAU [▲, Si/Al=15, space velocity 0.0074-0.03 mol

butanal (mol H+ s)

−1], and H4SiW12O40 [●, 0.075 mmol H4SiW12O40 gSiO2

−1, space velocity


−1] at 573 K. ........................................................................... 63

Figure 3.2. Carbon distributions of aromatic fraction produced in butanal reactions on (a-b) H-

FAU with different space velocities, (c) H-MFI, and (d) H4SiW12O40 at 573 K at time-on-stream

of 125 min. The distributions include aromatic molecules that do not lose any H ( ) or lose 2

( ), 4 ( ), or 6 ( ) hydrogen atoms in dehydrogenation reactions (e.g., Steps 1.1.2 and 1.3.3,

Scheme 3.1). .................................................................................................................................. 64

Figure 3.3. Rate ratios [ ,C H O,C H O 4 84 8

1-tetralin

( )j jr r , ,C H O,C H O 4 84 8

1-tetralin-ad

( )j jr r , or

,C H O,C H O 4 84 8

1-chd

( )j jr r ] for the rates of butanal primary reactions in (a) C4H8O-tetralin

(,C H O4 8 -tetralinjr ), (b) C4H8O-tetralin-adamantane (

,C H O4 8 -tetralin-adjr ), or (c) C4H8O-cyclohexadiene

(,C H O4 8 -chdjr ) feed mixtures to those in C4H8O feed ( ,C H O4 8jr ) as a function of tetralin pressure

( tetralinP ) or cyclohexadiene pressure ( chd

P ) for intermolecular C=C bond formation (Pathway 1,

♦), intramolecular C=C bond formation (Pathway 2, ●), isomerization-dehydration (Pathway 3,

▲), and Tishchenko esterification-ketonization (Pathway 4, ■) on H-FAU at 573 K [subscript

j=Inter, Intra, Dehy, or Tish, which denote inter- or intramolecular C=C bond formation,

isomerization-dehydration, or Tishchenko esterification-ketonization, respectively; space

velocity 0.0074 mol butanal (mol H+ s)

-1, adamantane (if added)=4-8 Pa]. The , j m values

(j=Inter, Intra, Dehy, or Tish; m=tetralin, tetralin-ad, or chd) are determined from the slopes in

xiv

these figures by linear regression of the data points against either Equations 3.5a, 3.5b, or 3.5c,

and are summarized in Table 3.2. ................................................................................................. 68

Figure 3.4. Rate constants for tetralin-to-alkanal (CnH2nO) transfer hydrogenation

(2,C H OTH -tetralinn n

k , Eqn. 3.9, n=3-6) on H-FAU as a function of the hydride ion affinity difference

between the protonated alkanal (CnH2nOH+) and the carbenium ion of tetralin (C10H11

+) (the

hydride ion affinity difference is given by Eqn. 3.7) [573 K, 1,1 kPa alkanal, 0.08-0.16 kPa

tetralin, space velocity 0.0074 mol alkanal (mol H+ s)

−1]. ........................................................... 79

Figure 4.1. (a) Rates (D4 8 2TH,C H O R H-r ) for butanal (C4H8O) transfer hydrogenation as a function of

H-donor pressure (D 2R HP ); (b)-(d) rate constants (

D4 8 2TH,C H O R H-k , Eqn. 4.3) for (b) butanal, (c)

propanal, and (d) pentanal transfer hydrogenation by various H-donors (RDH2), plotted as a

function of the hydride ion affinity difference ( + +D 2R H C H OH- n n

HIA , Eqn. 4.4) between the

carbenium ions of H-donor (RDH+, e.g., RDH

+=C10H11

+ for tetralin as the hydrogen donor) and

the protonated aldehydes (CnH2nOH+). The identities of H-donor are shown in the figure (573 K,

RDH2=cyclohexadiene, tetralin, cyclohexene, 3-methyl-1-pentene, or cyclohexane, H-FAU

(Si/Al=15)). The dash lines in (b)-(d) reflect the predicted reactivity trend of C6 H-donors

(cyclohexadiene, cyclohexene, and 3-methyl-1-pentene, which have similar molecular sizes)

based on + +D 2R H C H OH- n n

HIA . ....................................................................................................... 110

Figure 4.2. (a) and (b) Time-resolved infrared spectra upon exposure of H-FAU (Si/Al=15) to 10

Pa butanal followed by purging (a) in He or (b) in 15 Pa tetralin at 373 K; (c) coverages of

butanal on the H+ sites of H-FAU (

4 8C H O ) as a function of time upon purging with He or

introducing various H-donors at 373 K (line: fitted profiles against Eqn. 4.6); (d) butanal transfer

hydrogenation rate constants derived from in-situ infrared absorption spectroscopy (D 2TH-IR,R Hk ,

373 K) plotted against the transfer hydrogenation rate constants measured with steady-state

micro-catalytic flow reactor (D4 8 2TH,C H O R H-k , 573 K) with various H-donors (H-donors and

+ +D 4 8R H C H OH-

HIA values are indicated in the figures) ................................................................ 116

xv

Figure 4.3. The rate constants for aldehyde (CnH2nO, n=3-6) transfer hydrogenation with

cyclohexadiene (C6H8, 2 6 8TH,C H O C H-n n

k , , transition state depicted in Fig. 4.3(i)) or tetralin

(C10H12, 2 10 12TH,C H O C H-n n

k , , from [3], transition state depicted in Fig. 4.3(ii)) as the H-donor on

H-FAU zeolite (573 K), plotted as a function of hydride ion affinity difference

( + +10 11 2C H C H OH- n n

HIA or + +76 2C H C H OH- n n

HIA ); the rate constant for aldehyde transfer hydrogenation

with aromatic products as the H-donors (2,C H OTH n n

k , ▲, from[10], transition state depicted in Fig.

4.3(iii)) on H4SiW12O40 (573 K), plotted as a function of the hydride ion affinity difference

( + +10 11 2C H C H OH- n n

HIA ). .................................................................................................................. 119

Figure 5.1. (a) Overall butanal (C4H8O) conversion rates (◊) and carbon selectivities for C4H6

(○), C4H8 (∆), C8H14O (▼), C12H20O (■) , and C8+ hydrocarbons (labeled C8+ HC, ●) as a

function of time-on-stream during butanal reactions on H4SiW12O40 clusters at 573 K [butanal

pressure 1.1 kPa, 0.045 +1

butanal Hmol (mol s) , butanal conversion=18-24 %]; (b) H

+ site density,

expressed as the number of H+ site per H4SiW12O40 cluster remained after butanal reactions at

573 K, as a function of time-on-stream [butanal pressure 1.1-4.4 kPa, space velocity 0.045-0.18

+1

butanal Hmol (mol s) ]. ............................................................................................................... 142

Figure 5.2. Butanal conversions and carbon selectivities to (a) C8H14O (▼), C12H20O (■),

C16H26O (▲), and C8+ hydrocarbons [●, labeled C8+ HC, including C4tH6t aromatics (t=3 or 4),

cycloalkadienes (t=2), and C4tH6t+2 cycloalkenes (t=2)] and (b) C4H6 (○) and C4H8 (∆) during

butanal (C4H8O) reactions on H4SiW12O40 clusters [0.075 4 12 40 2

1H SiW O SiO

mmol g ] as a function

of space velocity at 623 K (1.1 kPa butanal in He, time-on-stream >155 min, at which stable

conversions and selectivities were attained). .............................................................................. 143

Figure 5.3 (a-d). Carbon distributions of the products, including oxygenates ( , from Steps 2a-

2c, 3a-3c, 4a, etc. in Scheme 5.1), aromatics ( , from Steps 3d-3e, etc.), cycloalkadienes ( ,

from Step 2d), n-dienes ( , from Step 1b), and n-alkenes ( , from Step 1a), during (a)

propanal, (b) butanal, (c) pentanal, and (d) hexanal reactions on H4SiW12O40 clusters at 573 K

[0.075 4 12 40 2

1H SiW O SiO

mmol g , space velocity=0.045 +

1alkanal H

mol (mol s) , alkanal pressure=1.1

3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 200

5

10

15

75

80

85

C5n

C4n

C3n

C2n

0.8%1%

0.7%0.6%2%

14%

1%

80%

Carb

on d

istr

ibution (

%)

Carbon number

Oxygenates

Aromatics

Cyclo-dienes

n-Dienes

n-Alkenes

0.4%

Cn

3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 200

5

10

15

75

80

85

C5n

C4n

C3n

C2n

0.8%1%

0.7%0.6%2%

14%

1%

80%

Carb

on d

istr

ibution (

%)

Carbon number

Oxygenates

Aromatics

Cyclo-dienes

n-Dienes

n-Alkenes

0.4%

Cn

3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 200

5

10

15

75

80

85

C5n

C4n

C3n

C2n

0.8%1%

0.7%0.6%2%

14%

1%

80%

Carb

on d

istr

ibution (

%)

Carbon number

Oxygenates

Aromatics

Cyclo-dienes

n-Dienes

n-Alkenes

0.4%

Cn

3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 200

5

10

15

75

80

85

C5n

C4n

C3n

C2n

0.8%1%

0.7%0.6%2%

14%

1%

80%

Ca

rbo

n d

istr

ibutio

n (

%)

Carbon number

Oxygenates

Aromatics

Cyclo-dienes

n-Dienes

n-Alkenes

0.4%

Cn

3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 200

5

10

15

75

80

85

C5n

C4n

C3n

C2n

0.8%1%

0.7%0.6%2%

14%

1%

80%

Carb

on d

istr

ibution (

%)

Carbon number

Oxygenates

Aromatics

Cyclo-dienes

n-Dienes

n-Alkenes

0.4%

Cn

xvi

kPa, time-on-stream=275 min, conversion=17 %, 30 %, 47 %, and 68 % for propanal, butanal,

pentanal, and hexanal, respectively]. .......................................................................................... 144

Figure 5.4. (a) Rates for intramolecular C=C bond formation ( Intrar , Cycle 2 in Scheme 5.2) as a

function of the total pressure of C8-C16 cyclic hydrocarbon product fraction ( HCP , including

cycloalkadienes and aromatics), (b) rates for intermolecular C=C bond formation ( Interr , Cycle 1

in Scheme 5.2) as a function of alkanal pressure ( alkanalP , average alkanal pressure), and (c) rates

for isomerization-dehydration ( Dehyr , Cycle 3 and Cycle 3.1 in Scheme 5.2) as a function of

alkanal pressure ( alkanalP , average alkanal pressure) during the reactions of alkanals [CnH2nO,

n=3-6; propanal (▲), butanal (■), pentanal ( ), and hexanal ( )] on H4SiW12O40 clusters [573

K, 0.045-0.44 +1

alkanal Hmol (mol s) , time-on-stream=275-600 min, alkanal conversion=14-17 %,

26-31 %, 45-47%, and 68-72 % for propanal, butanal, pentanal, and hexanal, respectively] .... 150

Figure 5.5. Rates for intermolecular C=C bond formation ( Interr , ■, Cycle 1 in Scheme 5.2) and

isomerization-dehydration ( Dehyr , ▲, Cycle 3 and Cycle 3.1 in Scheme 5.2) and the combined

rate ( Dehyr + Interr , ○ ) as a function of time-on-stream on H4SiW12O40 catalysts (0.075

4 12 40 2

1H SiW O SiO

mmol g ) at 573 K [space velocity=0.063 +

1butanal H

mol (mol s) , butanal

conversion=18-24 %]. ................................................................................................................. 152

Figure 5.6. (a) Rate constants for intermolecular C=C bond formation ( Inter,effk , , Cycle 1 in

Scheme 5.2), intramolecular C=C bond formation ( Intra,effk , ● , Cycle 2), and alkanal

isomerization-dehydration via bimolecular pathway ( Dehy,bi,effk , , Cycle 3) and monomolecular

pathway ( Dehy,mono,effk , ▲ , Cycle 3.1) and cyclization-dehydration selectivity of C3n alkenal

(3Cycli-dehy,C n

, ♦, Cyclization 2 and 2.1 in Scheme 5.2, Eqn. 5.15) during alkanal (CnH2nO, n=3-6)

deoxygenation on H4SiW12O40 clusters at 573 K as a function of reactant carbon number

[CnH2nO=1.1-10 kPa, 0.045-0.44 +1

alkanal Hmol (mol s) , alkanal conversion=14-17 %, 26-31 %,

45-47%, and 68-72 % for propanal, butanal, pentanal and hexanal, respectively]; (b) first-order

xvii

rate constants for intramolecular C=C formation ( Intra,effk ) for C3-C6 alkanals (CnH2nO, n=3-6) on

H4SiW12O40 clusters at 573 K as a function of the hydride ion affinity difference ( HIA )

between the carbenium ions of H-donor (R’H+) and the protonated alkanal (CnH2nOH

+) [ HIA =

+R'HHIA − +

2C H OHn nHIA , where R’H

+=C10H11

+, as tetralin (C10H12) is used as the representative H-

donor to estimate HIA values]. ................................................................................................ 167

Figure S2.1. Evolution of carbon in the products (○, excluding C3H6O) and unreacted C3H6O

(□) in the reactor effluent stream, total carbon in the reactor effluent stream (▲), and total

carbon in the feed mixture (dash line) as a function of time-on-stream during propanal (C3H6O)

reactions on H-MFI at 673 K (Si/Al=11.5, 5.3×10−2

mol C3H6O·(mol H+

i·s)−1

, 1.9 kPa C3H6O in

He). ................................................................................................................................................ 47

Figure S2.2. Propanal (C3H6O) conversion (■) and selectivities towards intermolecular C=C

bond formation (Sinter, ▲) and intramolecular C=C bond formation (Sintra, ●) as a function of

time-on-stream during propanal reactions on H-MFI at 473 K (Si/Al=11.5, 1.1×10−3

mol

C3H6O·(mol H+

i·s)−1

, 1.1 kPa C3H6O in He). ............................................................................... 48

Figure S2.3. Effects of H2O on rinter during propanal (C3H6O) reactions on the H-MFI

(Si/Al=11.5) at 473 K (7.5 ks, 1.1×10−3

mol C3H6O·(mol H+

i·s)−1

, 1.1 kPa C3H6O in He). ........ 49

Figure S3.1. (a) The rates for intermolecular C=C bond formation ( ,C H O4 8Interr , ♦) and Tishchenko

esterification-ketonization ( ,C H O4 8Tishr , □) as a function of H

+ site density and (b) the rates for

Tishchenko esterification-ketonization ( ,C H O4 8Tishr , □) as a function of basic site density during

butanal reaction on H-MFI and NaH-MFI zeolites at 573 K [1.1 kPa butanal, space

velocity=0.0037 mmol butanal (gcat. s) −1

, time-on-stream=125 min]........................................... 90

Figure S3.2. Concentration of remaining H+ sites on (a) H-MFI, (b) H-FAU, and (c) H4SiW12O40

catalysts after exposure to butanal reactants for different reaction times [573 K, space

velocity=0.0037-0.015, 0.0037, and 0.0074-0.030 mmol butanal (gcat. s)-1

for H-MFI, H-FAU,

and H4SiW12O40, respectively]. ..................................................................................................... 94

xviii

Figure S3.3. Butanal conversion rates (4 8,C H Ooverallr , ◊) and selectivities to intermolecular C=C

bond formation (4 8,C H OInterS , ● ), intramolecular C=C bond formation (

4 8,C H OIntraS , ▲ ),

isomerization-dehydration (4 8,C H ODehyS , ▼), and Tishchenko esterification-ketonization (

4 8,C H OTishS ,

■) during butanal reactions on (a) H-MFI, (b) H-FAU, and (c) H4SiW12O40 at 573 K as a

function of time-on-stream [1.1 kPa butanal, space velocity=0.0033, 0.0074, and 0.045 mol

butanal (mol H+ s)

−1 for H-MFI, H-FAU, and H4SiW12O40, respectively]. ................................. 95

Figure S3.4. Parity plot for the predicted and measured rates for intramolecular C=C bond

formation during butanal (C4H8O) reactions on H-MFI [ , space velocity 0.0033-0.013 mol

butanal (mol H+ s)

–1], H-FAU [▲, space velocity 0.0074-0.030 mol butanal (mol H

+ s)

−1], and

H4SiW12O40 [○, space velocity 0.045-0.18 mol butanal (mol H+ s)

−1] at 573 K. ....................... 96

Figure S3.5. Parity plots for the predicted and measured rate ratios [(a) ,C H O,C H O 4 84 8

1-tetralin

( )j jr r ,

(b) ,C H O,C H O 4 84 8

1-tetralin-ad

( )j jr r , or (c) ,C H O,C H O 4 84 8

1-chd

( )j jr r ] for rates of butanal reactions in (a)

C4H8O-tetralin (,C H O4 8 -tetralinjr ), (b) C4H8O-tetralin-adamantane (

,C H O4 8 -tetralin-adjr ), or (c) C4H8O-

cyclohexadiene (,C H O4 8 -chdjr ) feed mixtures to those in C4H8O feed ( ,C H O4 8jr ) for intermolecular

C=C bond formation (Pathway 1, ), intramolecular C=C bond formation (Pathway 2, ○),

isomerization-dehydration (Pathway 3, ∆), and Tishchenko esterification-ketonization (Pathway

4, ■) on H-FAU at 573 K [subscript j=Inter, Intra, Dehy, or Tish, which denote inter- or

intramolecular C=C bond formation, isomerization-dehydration, or Tishchenko esterification-

ketonization, respectively; space velocity 0.0074 mol butanal (mol H+ s)

-1, adamantane (if

added)=4-8 Pa]. ............................................................................................................................. 97


1-tetralin

( )j jr r

or (b) ,C H O,C H O 4 84 8

1-chd

( )j jr r ] for rates of butanal reactions in (a) C4H8O-tetralin (

,C H O4 8 -tetralinjr )

or (b) C4H8O-cyclohexadiene (,C H O4 8 -chdjr ) feed mixtures to those in C4H8O feed ( ,C H O4 8jr ) for

intermolecular C=C bond formation (Pathway 1, ), intramolecular C=C bond formation

(Pathway 2, ○ ), isomerization-dehydration (Pathway 3, ∆), and Tishchenko esterification-

0.00 0.05 0.10 0.15 0.200

1

2

3

4

rintra,with: rintra,without


rdehy,with:rdehy,without


rinter,with: rinter,without

rC7H14,with: rC7H14,without



rC7H14,with: rC7H14,withoutr C4H

8O

-C10H

12/r

C4H

8O

Tetralin pressure (kPa)

0.00 0.05 0.10 0.15 0.200

1

2

3

4










8O

-C10H

12/r

C4H

8O


0.00 0.05 0.10 0.15 0.200

1

2

3

4










8O

-C10H

12/r

C4H

8O


xix

ketonization (Pathway 4, ■) on H-MFI at 573 K [subscript j=Inter, Intra, Dehy, or Tish, which

denote inter- or intramolecular C=C bond formation, isomerization-dehydration, or Tishchenko

esterification-ketonization, respectively; space velocity 0.0033 mol butanal (mol H+ s)

-1]. ........ 98


1-tetralin

( )j jr r

or (b) ,C H O,C H O 4 84 8

1-chd





(Pathway 2, ○ ), isomerization-dehydration (Pathway 3, ∆), and Tishchenko esterification-

ketonization (Pathway 4, ■) on H4SiW12O40 at 573 K [subscript j=Inter, Intra, Dehy, or Tish,

which denote inter- or intramolecular C=C bond formation, isomerization-dehydration, or

Tishchenko esterification-ketonization, respectively; space velocity 0.045 mol butanal (mol H+

s)-1

]. ............................................................................................................................................... 99

Figure S3.8. Parity plots for the predicted and measured rates for tetralin-to-alkanal transfer

hydrogenation (,C H O2TH -tetralinn n

r , Eqn. 3.9, n=3-6) during CnH2nO-tetralin reactions (▲, □, ,

and for n=3, 4, 5, and 6, respectively) on H-FAU [573 K, space velocity 0.0074 mol CnH2nO

(mol H+ s)

−1, 1.1 kPa CnH2nO, 0.08-0.16 kPa tetralin]. .............................................................. 100

Figure S3.9. Parity plots for the predicted and measured rates for cyclohexadiene-to-butanal

transfer hydrogenation (,C H O4 8TH -chd

r , Eqn. 3.9) on H-FAU (▲), H-MFI (■), and H4SiW12O40

(●) at 573 K [space velocity=0.0074, 0.0033, and 0.045 mol butanal (mol H+ s)

−1 for H-FAU,

H-MFI, and H4SiW12O40, respectively, 1.1 kPa butanal, 0.03-0.15 kPa cyclohexadiene]. ........ 101

Figure S4.1. (a) Infrared spectra for H-FAU in He at 473 K (i), H-FAU exposed to 0.5 kPa

pyridine at 473 K (ii), and H-FAU exposed to 0.5 kPa butanal at 373 K (iii); (b) infrared spectra

for H-FAU exposed to 0.5 kPa butanal at 373 K. ....................................................................... 123

Figure S 4.2. Carbon distributions of aromatic fraction produced in butanal reactions on H-FAU

at 573 K (time-on-stream 125 min). The distributions include aromatic molecules that do not lose

any H ( ) or lose 2 ( ), 4 ( ), or 6 ( ) hydrogen atoms in dehydrogenation reactions.

0.00 0.05 0.10 0.15 0.200

1

2

3

4










8O

-C10H

12/r

C4H

8O


0.0 0.2 0.4 0.6 0.8 1.0 1.20.0

0.2

0.4

0.6

0.8

1.0

r TH

,Cn

H2

n-t

etr

alin

(10

-3 m

ol (m

ol H

+ s

)-1)

PTetralin (kPa)

0.0 0.2 0.4 0.6 0.8 1.0 1.20.0

0.2

0.4

0.6

0.8

1.0

r TH

,Cn

H2

n-t

etr

alin(1

0-3

mo

l (m

ol H

+ s

)-1)

PTetralin (kPa)

xx

Examples of the C12 aromatic products are given. These results have been published elsewhere

[3]. ............................................................................................................................................... 124

Figure S4.3. (a) Differential infrared spectra of H-FAU during butanal adsorption and desorption

by subtracting the spectrum of pristine H-FAU; (b) the total H+ site coverage during butanal

adsorption and desorption on H-FAU as a function of time (373 K, butanal pressure 10 Pa,

helium flow rate 0.83 cm3 s

−1). ................................................................................................... 126

Figure S4.4. (a) Infrared spectra of butanal adsorption on H-MFI zeolite (Si/Al=40) at 308 K in

0.01 Pa butanal; (b) correlation between the decrease of band area for the H+ site (−∆Av(OH), MFI,

stretching vibration at 3610 cm−1

) and the increase of the band area for the adsorbed carbonyl

group (∆Av(C=O), stretching vibration at 1670 cm−1

) during butanal adsorption on H-MFI zeolite

(Si/Al=40) at 308 K in 0.01 Pa butanal. ...................................................................................... 129

Figure S4.5. Time-resolved infrared spectra of H-FAU zeolite upon butanal adsorption (10 Pa

butanal, gray lines) followed by purging in (a) cyclohexadiene (10 Pa), (b) cyclohexene (11 Pa),

and (c) cyclohexane (12 Pa) at 373 K. ........................................................................................ 130

Figure S5.1. H+ site densities on H3PW12O40 clusters after thermal treatment under flowing He at

different temperatures (473-677 K) for 180 min (loading amount 0.13 12 40 2

1H PW O SiO

mmol g ). 178

Figure S5.2. The differential spectra of H4SiW12O40 before and after butanal reactions at 573 K

for 5 min. ..................................................................................................................................... 179

Figure S5.3. Weight, differential weight loss, and heat flow profiles during the temperature

programmed oxidation (TPO) of spent H4SiW12O40/SiO2 catalysts (0.075 4 12 40 2

1H SiW O SiO

mmol g )

after butanal reactions (0.5 kPa butanal) at 573 K for 8 h (initial sample weight=15.3 mg, in 5%

O2/He, 10 cm3 min

−1, heating rate =5 K min

−1). ......................................................................... 180

Figure S5.4. Rates for (a) intermolecular C=C bond formation ( Interr ), (b) intramolecular C=C

bond formation ( Intrar ), and (c) isomerization-dehydration ( Dehyr ) during butanal reaction on H-

MFI zeolite (Si/Al=40) and the ammonium hexafluorosilicate treated H-MFI zeolite [H-

xxi

MFI(AHFS)] at 573 K as a function of time-on-stream [butanal=4 kPa, space velocity=1.5×10−5

mol (gcat. s)−1

]. ............................................................................................................................. 182

Figure S5.5. Profiles of pyridine-TPD for (a) fresh H4SiW12O40 catalysts and (b) spent

H4SiW12O40 catalysts after 12 h of butanal reactions at 573 K (pyridine adsorption temperature

473 K, heating rate 1 K min−1

). ................................................................................................... 183

Figure S5.6. Carbon selectivities of the products during 2,4-heptadienal reactions on H4SiW12O40

clusters at 573 K [2,4-heptadienal pressure=0.2 kPa, space velocity=0.009 +1

Hmol (mol s) ,

time-on-stream=125 min] ........................................................................................................... 184

xxii

List of Schemes

Scheme 2.1. Reaction network for propanal turnover on H-MFI zeolite (#: Intermediates at

undetectable concentrations; *: Hydrogen donating agents for deoxygenation reaction in R 2.2).

....................................................................................................................................................... 19

Scheme 2.2. Mechanism for inter- and intramolecular C=C bond formation in propanal (C3H6O)

evolving 2-methyl-2-pentenal (C6H10O) and propylene, respectively, on H+ sites ( denotes

quasi-equilibrated step, reversible step, and irreversible step). In taut,' sK of Step

2.1b, species s denotes propanal surface isomers [(4a), (4b), and their physisorbed isomers]. .... 25

Scheme 3.1. Reaction network for butanal deoxygenation on solid Brønsted acid catalysts (“D”

and “A” denote H-donor and H-acceptor, respectively; most of the intermediates and products

shown in the scheme were detected in the experiment except crotyl alcohol and butanol because

of their rapid dehydration). ........................................................................................................... 59

Scheme 3.2. A proposed mechanism for intermolecular hydride transfer from tetralin to

protonated butanal (the kinetically relevant step for Pathway 2). ................................................ 77

Scheme 3.3. Reaction network for CnH2n-tetralin (naphthalene) alkylation. ................................ 78

Scheme 4.1. Pathways of transfer hydrogenation of (a) protonated quinoline catalyzed by chiral

Brønsted acid (chiral phosphoric acid) [4], (b) pyruvate catalyzed by lactate dehydrogenase [8],

(c) protonated alkene catalyzed by solid Brønsted acid sites (e.g., H-MOR zeolite [5], and H3Si-

OH-AlH2-O-SiH3 cluster model [7]), and (d) protonated aldehyde catalyzed by solid Brønsted

acid sites (R, R1, and R2 denote alkyl groups). In each of these cases, the reaction involves

protonation of the reactant followed by hydride transfer (from a hydride donor) (H-acceptor and

H-donor denote hydride acceptor and hydride donor, respectively). ......................................... 104

Scheme 4.2. Pathways of aldehyde reactions that generate hydrogen donors (aromatics or their

precursors, labeled H-donor), which include aldol condensation and ring closure steps (R and R1-

R4 represent either an H or alkyl groups). The parallel pathways of aromatic transalkylation and

their products are not shown in the scheme for simplification purposes (the complete reaction

network is reported elsewhere[3]). ............................................................................................. 108

xxiii

Scheme 4.3. (a) Catalytic steps (Steps 1-3) for solid Brønsted acid catalyzed transfer

hydrogenation of aldehydes (CnH2nO) by hydrogen donors (RDH2) and the proposed hydride

transfer transition state; (b) the heat of reaction for the kinetically-relevant hydride transfer step

( THrH , Step 1), interpreted using a Born-Haber thermochemical construct ( +DR H

HIA and

+2C H OHn n

HIA denote the hydride ion affinities of the H-donor carbenium ion and protonated

aldehyde, respectively). .............................................................................................................. 112

Scheme 5.1. Pathways for alkanal (CnH2nO) chain growth resulting in larger alkenals (and their

isomers, CtnH2tn-2t+2O, n=3-6, t=2-3) and hydrocarbons (including cycloalkadienes and aromatics,

CtnH2tn-2t, n=3-6, t=2-3) (R, R1, and R2 represent either alkyl group or H)................................. 146

Scheme 5.2. Reaction network for alkanal (CnH2nO, n=3-6) deoxygenation on H4SiW12O40

clusters capturing the intermolecular C=C bond formation (Cycle 1, 1.1, and 1.2), intramolecular

C=C bond formation (Cycle 2), isomerization-dehydration (Cycle 3 and 3.1), the secondary

cyclization-dehydration and dehydrogenation reactions (Cyclization 1, 2, 2.1, and 3), illustrated

with butanal as an example. ........................................................................................................ 148

Scheme 5.3. Mechanism for intermolecular C=C bond formation (Steps G1, A1, and R1.1-R1.5,

also shown as Cycle 1 in Scheme 5.2), intramolecular C=C bond formation (Steps A1 and R2.1-

R2.4, also shown as Cycle 2 in Scheme 5.2), and isomerization-dehydration via bimolecular

pathway (Steps R3.1a-R3.2a, also shown as Cycle 3 in Scheme 5.2) and monomolecular pathway

(Steps R3.1b-R3.2b, also shown as Cycle 3.1 in Scheme 5.2) during alkanal reactions on

H4SiW12O40 clusters (R=H, CH3, C2H5, and C3H7 for propanal, butanal, pentanal, and hexanal,

respectively; R’H2 represents a H-donor). .................................................................................. 153

Scheme 5.4. (a) Proposed mechanism for acid catalyzed cyclization-dehydration of 2,4-

heptadienal (C7H10O) [the products detected are labeled with carbon selectivities within the C7

product fractions during 2,4-heptadienal reactions on H4SiW12O40 at 573 K, 2,4-heptadienal

pressure 0.2 kPa, space velocity= 0.009 +1

Hmol (mol s) , time-on-stream=125 min]; (b)

proposed mechanism for acid catalyzed cyclization-dehydration of 2,4-diethyl-2,4-octadienal

(C12H20O) during butanal reactions [the products detected are labeled with carbon selectivities

xxiv

within the C12 product fractions during butanal reaction on H4SiW12O40 at 573 K, butanal

pressure 1.1 kPa, space velocity=0.045 +1

Hmol (mol s) , time-on-stream=125 min]. .............. 164

Scheme S3.1. Thermochemical cycles used for estimating the hydride ion affinities (HIA) for (a)

the carbenium ion (R’H+) of a hydrocarbon (R’H2, taking R’H2=tetralin as an example) and (b)

the protonated alkanal (CnH2nOH+) [ R'PA and

2C H On nPA are the proton affinities of hydrocarbon

R’ and alkanal CnH2nO, respectively; r H ionH (−1675.3 kJ mol

-1) [47] is the heat of reaction

for H++H

− → H2; r Hydro,R'

H and 2

r Hydro,C H On nH are the heats of reaction for hydrogenation

reactions: R’+H2→ R’H2 and CnH2nO+ H2→CnH2n+1OH, respectively]. .................................. 92


the carbenium ion (RDH+) of a hydrocarbon (RDH2, taking RDH2=tetralin as an example) and (b)

the protonated aldehyde (CnH2nOH+) [

DRPA and 2C H On n

PA are the proton affinities of

hydrocarbon RD and aldehyde CnH2nO, respectively; r H ionH (−1675.3 kJ mol-1

) [25] is the heat

of reaction for H++H

- → H2;

Dr Hydro, RH and 2, C H Or Hydro n n

H are the heats of reaction for

hydrogenation reactions: RD+H2→ RDH2 and CnH2nO+ H2→CnH2n+1OH, respectively]........... 132

Preface

This thesis is based on manuscripts that have been published in or submitted for publication in

peer reviewed journals. Consequently, there may be some overlap in material that is presented

throughout the thesis. All manuscripts included in this thesis were written by Fan Lin, with

critical comments provided by Dr. Ya-Huei (Cathy) Chin. Contributions of any other people are

described below:

Chapter 2

Published as: F. Lin, Y.-H. Chin, “Mechanism of intra- and inter-molecular C=C bond formation

of propanal on Brønsted acid sites contained within MFI zeolites”, J. Catal., 311 (2014) 244-256.

Contributions: The experimental approach was developed by Fan Lin and Ya-Huei (Cathy) Chin.

The infrared measurements were performed by Yuanshuai Liu and Eszter Barath. The sodium

exchanged zeolite samples were prepared by Han-Yue Fu. All the other experiments described in

this section were performed by Fan Lin. The manuscript was written by Fan Lin with critical

comments from Ya-Huei (Cathy) Chin.

Chapter 3

Published as: F. Lin, Y.-H. Chin, “Alkanal Transfer Hydrogenation Catalyzed by Solid Brønsted

Acid Sites”, J. Catal., 341(2016) 136-148.


The sodium exchanged zeolite samples were prepared by Han-Yue Fu. All the other experiments

described in this section were performed by Fan Lin. The manuscript was written by Fan Lin

with critical comments from Ya-Huei (Cathy) Chin.

Chapter 4

To be submitted as: F. Lin, Y. Yang, Y.-H.C. Chin, “Kinetic Requirements of Solid Brønsted

Acid Catalyzed Transfer Hydrogenations of Aldehyde”, Angew. Chem. Int. Ed.


The infrared measurements were performed by Yifei Yang. All the other experiments described

xxvi

in this section were performed by Fan Lin. The manuscript was written by Fan Lin and Yifei

Yang, with critical comments from Ya-Huei (Cathy) Chin.

Chapter 5

Published as: F. Lin, Y.-H.C. Chin, “Catalytic Pathways and Kinetic Requirements for Alkanal

Deoxygenation on Solid Tungstosilicic Acid Clusters”, ACS Catal., 6 (2016), 6634-6650.


The infrared measurements were performed by Yifei Yang. All the other experiments described

in this section were performed by Fan Lin. The manuscript was written by Fan Lin with critical

comments from Ya-Huei (Cathy) Chin.

1

Chapter 1 Introduction to Alkanal Deoxygenation on Solid Brønsted Acid

Catalysts

Catalytic deoxygenation of alkanals (or aldehydes) to hydrocarbons on solid Brønsted acid

catalysts (e.g. acidic zeolites), under ambient pressure and moderate temperatures (473-723 K),

is a potential route to upgrade the alkanal fragments in the biomass pyrolysis oils for the

production of drop-in liquid fuels and value-added chemicals. There are multiple reaction routes

for solid Brønsted acid catalyzed alkanal deoxygenation, leading to the formation of light alkenes,

alkadienes, and heavy alkenals and aromatics. In this thesis, kinetic studies were carried out to

gain the mechanistic insights into these deoxygenation pathways, and to uncover their catalytic

requirements. The reactivity trend across the alkanal family with different molecular sizes and

the kinetic consequences of the local confinement of the active sites were also investigated.

Fast pyrolysis of lignocellulosic biomass produces light carboxylic acids (RCOOH), aldehydes

(RCHO), ketones [R(C=O)R’], and alcohols (ROH) with less than or equal to six carbon atoms

(R and R’ represent alkyl groups) [1, 2]. As an example, pyrolysis of bark free wood at short

residence times (1-2 s) and 793 K leads to organic liquid fractions containing: (1) 5-10 wt.%

formic acid, methanol, and formaldehyde (C1), (2) 15-35 wt.% linear hydroxyl and oxo-

substituted aldehydes and ketones with two to four carbon atoms (C2-C4), (3) 10-20 wt.%

hydroxyl, hydroxymethyl, and/or oxo-substituted furans, furanones, and pyranones with five to

six carbon atoms (C5-C6), and (4) 6-10 wt.% anhydrosugars including anhydro-oligosaccharides

(C6) [1]. Deoxygenation and carbon-carbon bond formation reactions convert these oxygenates,

by removal of their oxygen atoms and augmenting their carbon chain length, on solid Brønsted

acid catalysts (e.g. H-ZSM-5 [3-6], H-MOR [5], H-FAU [5, 6] zeolites) at ambient pressure and

moderate temperatures (563-723 K) to hydrocarbons as a route to produce drop-in liquid fuels

(C6-C20 hydrocarbons) and value-added chemicals (e.g., alkadienes).

Contained within the light oxygenate fraction are alkanals (mostly C2-C6, e.g.,

hydroxyacetaldehyde and furfural, ~20 wt.% of the organic fraction [2, 7]). The reactions of

alkanal species on solid Brønsted acid sites have been probed using C3-C6 n-alkanals (CnH2nO,

n=3-6) as model reactants [8-12], and three concomitant primary pathways were identified: (i)

intermolecular C=C bond formation, a bimolecular pathway lengthening the carbon chain via

2

aldol condensation and dehydration and forming larger alkenals (C2nH4n-2O), (ii) intramolecular

C=C bond formation, an unimolecular pathway evolving alkene (CnH2n) via transfer

hydrogenation and dehydration while preserving the carbon backbone, and (iii) isomerization-

dehydration, another unimolecular pathway directly ejecting a water molecule from the alkanal

and producing alkadiene (CnH2n-2). The alkenal product of pathway (i) can undergo secondary

intermolecular C=C formation with another alkanal and further lengthen the carbon chain,

followed by cyclization-dehydration, leading to the formation of aromatic products. These

reaction routes for alkanal deoxygenation on H+ sites have been reported individually in previous

studies, however, with few details about the reaction mechanisms and kinetic requirements,

except for the mechanism of the acid-catalyzed aldol condensation [13, 14].

The catalytic deoxygenation and C=C bond coupling of alkanals via aldol condensation and the

following dehydration have been reported on solid acid catalysts (e.g. H-MFI [4, 8, 9]). The

mechanisms for solid acid-catalyzed aldol condensation have been well established (e.g. on H-

MFI [13] and H-Y [14] zeolites). This reaction is initiated by keto-enol tautomerization of

alkanal to form a small concentration of alkenol in the gas phase, which then nucleophilically

attacks the carbonyl carbon of the protonated alkanal adsorbed on the Brønsted acid site (H+),

creating an intermolecular C-C bond and forming an aldol compound. The aldol rapidly

dehydrates on the acid sites and leading to an alkenal with a lengthened carbon chain. The

sequential condensation reactions further augment the carbon numbers, before the eventual

intramolecular carbon-carbon bond formation that closes the carbon ring, followed by

dehydration, dehydrogenation, and transalkylation to evolve diverse aromatics. For example,

deoxygenation of propanal (C3H6O) on H-ZSM-5 zeolites at 673 K forms predominantly C6-C10+

aromatics (carbon selectivity 42-53 %) [8, 9]. Although the mechanism of the aldol condensation

step has been established, few studies have addressed the secondary reactions that lead to the

formation of larger olefinic or aromatic products.

A separate reaction for alkene formation from alkanal also occur, as reported previously for

alkanal reactions on H-ZSM-5 zeolite [8]. Propanal reactions on H-ZSM-5 zeolite at 673 K

produced a significant amount of C1-C3 light gases (43-53% carbon selectivity) in which the

propene was the predominant product [8]. The light products including propene were proposed to

come from the cracking of larger aromatics, but without a detailed mechanistic description.

Reactions of CnH2nO alkanals (n=3-5) on H-ZSM-5 zeolite produce near exclusive CnH2n alkenes

3

(CnH2n/ 1

21C H

t n

t tt n

=0.93, 0.95, and 0.89 for n=3, 4, and 5, respectively, at 473 K) [10]. The

reaction stoichiometry indicates that the alkene formation from alkanal likely occurs via a direct

hydrogen transfer step, during which an alkanal accepts two hydrogen atoms, followed by

dehydration, leaving the carbon backbone untouched. The reaction pathway of hydrogen transfer

onto a carbonyl group followed by dehydration was also proposed in a recent work studying the

hydrogen transfer and dehydration of naphthols on H-Y zeolite [15]. An increase in hydrogen

transfer rates was observed in the presence of hydrocarbons (e.g., tetralin and 1,5-

dimethyltetralin), as these hydrocarbons acted as the hydrogen donors which shuffled hydride

ion to the keto tautomers of naphthols, and the rate of such events appeared to correlate to the

hydride ion dissociation energies of the H-donors. Nevertheless, to date, little mechanistic detail

is available for the transfer hydrogenation of n-alkanals, despite the clear kinetic evidence of

their predominant occurrences during their reactions on Brønsted solid acid catalysts.

Other than the pathway of intramolecular C=C bond formation, alkanals (CnH2nO) may remove

their oxygen via a direct isomerization-dehydration route, which forms the corresponding

alkadienes (CnH2n-2) [16-19]. In fact, previous studies have shown that 2-methylbutanal

dehydration on borosilicate zeolite [16] or aluminum phosphate (AlPO4) [18, 19] leads to

isoprene and 2-methylpentanal dehydration on aluminosilicate zeolite (H-Y) to 2-methylpenta-

1,3-diene [16], respectively, as viable routes for synthesizing polymer precursors. This alkanal

dehydration reaction was proposed to occur via allylic alcohol intermediates, because (1) 2-

methylbutanal and its ketone isomer (methyl isopropyl ketone) and allylic alcohol isomer (2-

methyl-2-buten-1-ol) can interconvert with each other (on BPO4 and AlPO4 catalysts, at 383 K),

and (2) both 2-methylbutanal and isopropyl ketone give similar yields to isoprene on AlPO4 (54 %

vs. 49 % at 673 K). Therefore, it was proposed that the reactions must involve a common allylic

alcohol intermediate, which connects the two isomers, 2-methylbutanal and methyl isopropyl

ketone, and the dehydration product, isoprene [19].

Multiple concomitant catalytic routes are evident from these previous studies for alkanal

deoxygenation on solid Brønsted acid catalysts. These catalytic routes result in heavy oxygenates

and aromatics, as well as light alkenes and alkadienes. The individual rates of these routes, their

kinetic requirements, and the kinetic connection between these pathways have, however,

remained largely unresolved. Questions about how the activities of these deoxygenation routes

4

vary with the alkanal molecular size and the Brønsted site local environment are of great interest,

as this knowledge would serve as guidelines for catalyst and process design to control the

product selectivity of the alkanal deoxygenation.

In Chapter 2, propanal was used as a model compound to study the alkanal deoxygenation on H+

sites contained within H-MFI zeolites at median temperature (473-673 K). The kinetic

measurements of propanal reactions on H-MFI zeolites, together with the H+ site characterization

(including pyridine titration, pyridine IR, and propanal TPD), were carried out to establish the

pathways of alkanal deoxygenation on H+ sites and their reaction kinetics. Specifically, propanal

deoxygenation proceeds via two primary pathways of inter- and intramolecular C=C bond

formation. In the pathway of intermolecular C=C bond formation, two propanal molecules

undergo aldol condensation-dehydration to form a larger alkenal (2-methyl-2-pentenal). This

pathway adds additional propanal units to further lengthen the carbon chain of the alkenals,

which then undergo cyclization-dehydration to produce aromatics. In parallel, the pathway of

intramolecular C=C bond formation involves hydrogen transfer from the aromatic products to a

single propanal molecule to form a propanol, which undergoes sequential dehydration to produce

propene while preserving the carbon backbone. Kinetic models of these two pathways are

established and validated with kinetic measurements under various conditions (varied partial

pressures of alkanal reactant, co-feed H2O and hydrogen donor (3-methyl-1-pentene), etc.).

The more detailed mechanism and kinetic requirement of the hydrogen transfer in the pathway of

intramolecular C=C bond formation (also called alkanal transfer hydrogenation) are probed in

Chapter 3, by varying the molecular size of the alkanal reactants (C3-C6 n-alkanals) and the local

structure of the H+ sites (either immobilized in microporous MFI and FAU crystalline structures

or dispersed on H4SiW12O40 polyoxometalate clusters). Here I establish the kinetic sensitivity of

the alkanal transfer hydrogenation on the thermochemical properties of the reactants and local

environment of H+ sites. Transfer hydrogenation of alkanals involves a kinetically relevant,

intermolecular hydride transfer step from aromatic or cycloalkadiene products (e.g. substituted

tetralins and cyclohexadienes), as hydride donors (R’H2), to protonated alkanals (RCH2CHOH+)

as the hydride acceptors, via a bimolecular transition state with a shared hydride ion:

(RCH2CHOH+∙∙∙H

−∙∙∙R’H

+)‡. The rate constants of transfer hydrogenation are determined by the

hydride ion affinity difference between the carbenium ions of the H-donors (R’H+) and the

protonated alkanals (RCH2CHOH+). The transfer hydrogenation occurs much more effectively

5

on partially confined H+ sites in FAU structures than in smaller pore MFI or unconfined

H4SiW12O40 polyoxometalate clusters, which is an indication that FAU solvates and stabilizes

the bulky transition state of hydride transfer via van der Waals interactions.

In Chapter 4, the mechanism and kinetic requirement for alkanal transfer hydrogenation were

further confirmed by varying the identities of the H-donors (hydrocarbon species) and H-

acceptors (C3-C6 alkanals), and by correlating the kinetic measurement results obtained via the

micro-catalytic reactor experiments and the in-situ infrared spectroscopic study, respectively.

Incorporating H-donors (e.g. cyclohexadiene, tetralin, and cyclohexene) into alkanal (CnH2nO,

n=3-6) reactions on H-FAU zeolites promotes the alkanal transfer hydrogenation and increases

the rate of alkene (CnH2n) formation. The in-situ infrared spectroscopic experiments show that

these H-donors also accelerated the disappearance of the protonated butanal on the H-FAU

zeolite by transfer hydrogenation. Both kinetic measurement methods give consistent reactivity

trends for four pairs of H-donor-acceptor, which confirms that the hydride ion affinity difference

between a H-donor-acceptor pair is the kinetic descriptor for the hydride transfer. In addition, the

size fitting between the transition state (RCH2CHOH+∙∙∙H

−∙∙∙R’H

+)‡ and the H

+ site local

structure plays a critical role in tuning the stability of the transition state and the rate constant.

The hydride transfer is most favorable when the transition state is solvated and stabilized in a

cage that matches the size of the transition state. If the transition state is too small and loosely

fitted in the cage, this solvation effect becomes weaker; on the other hand, if the transition state

is oversized in the cage, the steric constraint destabilizes the transition state. Both cases lead to

higher activation barriers and to lower hydride transfer rate constants.

Other than the two primary pathways of inter- and intramolecular C=C bond formation, there are

multiple catalytic routes occurring concomitantly during alkanal deoxygenation on solid

Brønsted acid sites, including another primary route of alkanal isomerization-dehydration

producing light alkadienes, the secondary reactions of the intermolecular C=C bond formation

further lengthening the carbon chain of alkenals, the alkenal cyclization-dehydration evolving

aromatic species, and the aromatic transalkylation leading to diverse-sized aromatics. The

complexity of the reaction system causes the ambiguity of the catalytic pathways and the

associated mechanisms. On the other hand, the relative rates of these steps are expected to vary

with the thermodynamic properties of the acid sites and their environments, such as acid strength

and site confinement afforded by pores and cages of molecular dimensions. Using catalysts with

6

diverse site environments (e.g., microporous crystalline materials, H-MFI and H-FAU zeolites)

in probing these inherently complex pathways has further made kinetic data interpretation and

their connection to the structural and thermodynamic properties of active sites difficult.

In Chapter 5, I probe the catalytic pathways of alkanal deoxygenation with strategies of kinetic

measurement and chemical titration, after isolating the kinetic contributions of acid strengths and

site environments. I focus on the deoxygenation chemistry of straight chain alkanals (n-CnH2nO)

with three to six carbon atoms, carried out on Keggin polyoxometalate clusters (tungstosilicic

acid, H4SiW12O40) with well-defined structures and isolated, uniformed H+ sites anchored on the

cluster surfaces and without the local molecular confinements typically found in microporous

crystalline materials. I probe the reaction pathways by systematically examining the primary and

secondary reactions, by quantitative kinetic studies decoupling the rate contributions from the

various routes. I construct a reaction network reconciling the various pathways and diverse

products reported across the literature, propose reaction mechanism, and correlate the

thermodynamic properties to the rates of the individual reaction routes of alkanal intermolecular

C=C bond formation, intramolecular C=C bond formation, and isomerization-dehydration. This

approach provides simple explanations of the complex reaction systems and correlates

thermodynamic properties such as site environment and alkanal chain lengths to rates and

selectivities during deoxygenation reaction.

The following chapters will provide the details on the kinetic measurements and active site

characterization which uncover the complex reactions of alkanal deoxygenation on solid

Brønsted acid sites. The mechanistic insights into the deoxygenation routes gained in this work

will unravel the catalytic and kinetic requirements of alkanal deoxygenation. The knowledge of

how the alkanal reactant molecular sizes, their thermodynamic properties, and the active site

local structures affect the activity of individual deoxygenation route allow catalyst and process

design to tune the reaction selectivities leading to the various products of alkenals, aromatics,

light alkenes and alkadienes.

7

References

[1] A. Oasmaa, E. Kuoppala, S. Gust, Y. Solantausta, Fast Pyrolysis of Forestry Residue. 1.

Effect of Extractives on Phase Separation of Pyrolysis Liquids, Energy Fuels, 17 (2003) 1-12.

[2] B. Valle, A.G. Gayubo, A.T. Aguayo, M. Olazar, J. Bilbao, Selective Production of

Aromatics by Crude Bio-oil Valorization with a Nickel-Modified HZSM-5 Zeolite Catalyst,

Energy Fuels, 24 (2010) 2060-2070.

[3] A.G. Gayubo, A.T. Aguayo, A. Atutxa, R. Aguado, J. Bilbao, Transformation of Oxygenate

Components of Biomass Pyrolysis Oil on a HZSM-5 Zeolite. I. Alcohols and Phenols, Ind. Eng.

Chem. Res., 43 (2004) 2610-2618.

[4] A.G. Gayubo, A.T. Aguayo, A. Atutxa, R. Aguado, M. Olazar, J. Bilbao, Transformation of

Oxygenate Components of Biomass Pyrolysis Oil on a HZSM-5 Zeolite. II. Aldehydes, Ketones,

and Acids, Ind. Eng. Chem. Res., 43 (2004) 2619-2626.

[5] J.D. Adjaye, N.N. Bakhshi, Production of hydrocarbons by catalytic upgrading of a fast

pyrolysis bio-oil. Part I: Conversion over various catalysts, Fuel Process. Technol., 45 (1995)

161-183.

[6] J. Jae, G.A. Tompsett, A.J. Foster, K.D. Hammond, S.M. Auerbach, R.F. Lobo, G.W. Huber,

Investigation into the shape selectivity of zeolite catalysts for biomass conversion, J. Catal., 279

(2011) 257-268.

[7] Q. Zhang, J. Chang, T. Wang, Y. Xu, Review of biomass pyrolysis oil properties and

upgrading research, Energ. Convers. Manage., 48 (2007) 87-92.

[8] T.Q. Hoang, X. Zhu, T. Sooknoi, D.E. Resasco, R.G. Mallinson, A comparison of the

reactivities of propanal and propylene on HZSM-5, J. Catal., 271 (2010) 201-208.

[9] T.Q. Hoang, X. Zhu, L.L. Lobban, D.E. Resasco, R.G. Mallinson, Effects of HZSM-5

crystallite size on stability and alkyl-aromatics product distribution from conversion of propanal,

Catal. Commun., 11 (2010) 977-981.

[10] F. Lin, Y.-H. Chin, Mechanism of intra- and inter-molecular C=C bond formation of

propanal on Brønsted acid sites contained within MFI zeolites, J. Catal., 311 (2014) 244-256.

[11] F. Lin, Y.-H. Chin, Alkanal Transfer Hydrogenation Catalyzed by Solid Brønsted Acid Sites,

J. Catal., 341 (2016) 136-148.

[12] F. Lin, Y.-H. Chin, Catalytic Pathways and Kinetic Requirements for Alkanal

Deoxygenation on Solid Tungstosilicic Acid Clusters, ACS Catal., 6 (2016) 6634-6650.

[13] E. Dumitriu, V. Hulea, I. Fechete, A. Auroux, J.-F. Lacaze, C. Guimon, The aldol

condensation of lower aldehydes over MFI zeolites with different acidic properties, Micropor.

Mesopor. Mat., 43 (2001) 341-359.

8

[14] T. Komatsu, M. Mitsuhashi, T. Yashima, Aldol Condensation Catalyzed by Acidic Zeolites,

in: R. Aiello, G. Giordano, F. Testa (Eds.) 2nd International Conference of the Federation-of-

European-Zeolite-Associations Taormina, Italy, 2002, p. 667-674.

[15] T. Prasomsri, R.E. Galiasso Tailleur, W.E. Alvarez, T. Sooknoi, D.E. Resasco, Conversion

of 1-tetralone over HY zeolite: An indicator of the extent of hydrogen transfer, Appl. Catal. A,

389 (2010) 140-146.

[16] W. Hoelderich, F. Merger, W.D. Mross, G. Fouquet, Preparation of Dienes by Dehydration

of Aldehydes, US4560822 A (1985).

[17] L.G. Wideman, T. Ohio, Process for the Production of Diene from Aldehydes, US4628140

(1986).

[18] I.D. Hudson, G.J. Hutchings, Preparation of Conjugated Dienes, US5264644 A (1993).

[19] G.J. Hutchings, I.D. Hudson, D. Bethell, D.G. Timms, Dehydration of 2-Methylbutanal and

Methyl Isopropyl Ketone to Isoprene Using Boron and Aluminium Phosphate Catalysts, J. Catal.,

188 (1999) 291-299.

9

Chapter 2 Mechanism of Intra- and Intermolecular C=C Bond Formation

of Propanal on Brønsted Acid Sites Contained within MFI Zeolites

Abstract

Kinetic and chemical titration studies are used to unravel the reaction pathways and catalytic

requirements for propanal deoxygenation over Brønsted acid sites contained within MFI zeolites.

Propanal deoxygenation in the absence of an external hydrogen source is initiated via primary

and competitive pathways of inter- and intramolecular C=C bond formation that involve

bimolecular coupling of propanal and unimolecular deoxygenation steps, respectively. The

intermolecular C=C bond formation proceeds via mechanistic steps resembling the acid

catalyzed aldol condensation reactions in the homogeneous phase and its reactive collision

frequencies increase with increasing propanal pressure. The reaction is initiated by keto-enol

tautomerization of propanal to form small concentrations of propenol. The propenol undergoes

kinetically-relevant nucleophilic attack to protonated propanal, the most abundant surface

intermediates, to create the intermolecular C=C bond. The competitive unimolecular

deoxygenation step involves kinetically relevant hydrogen transfer from hydrogen donating

agents and occurs at rates that remain invariant with propanal pressure. Hydrogen donating

agents are aliphatic rings produced from consecutive intermolecular C=C bond formation and

ring closure events, and donate hydrogen via dehydrogenation steps to increase their extent of

unsaturation. Hydrogen donating events must be kinetically coupled with the direct hydrogen

addition step on propanal to satisfy the deoxygenation stoichiometry and form propanol, which

upon dehydration evolves predominantly propene, thus preserving the carbon backbone. Water

as a byproduct prevents binding of larger, inactive carbonaceous species on acid sites and

inhibits the intermolecular C=C bond formation step by increasing the reverse rate of this step.

Water, however, does not alter the net rate for intramolecular C=C bond formation, because of its

irreversible nature. An increase in the rate ratio for intra- over intermolecular C=C bond

formation upon the addition of 3-methyl-1-pentene, an effective hydrogen donating agent,

confirms the kinetic relevance of the hydrogen transfer step for propene formation. These

findings on the different kinetic dependencies for the competitive reactions and their mechanistic

10

interpretations provide the operating strategies to tune the reaction pathways, manipulate the

extent of hydrogen transfer, and tailor the distributions of larger oxygenates and alkenes during

propanal deoxygenation reactions.

2.1. Introduction

Small oxygenates of alkanal and alkanone (R-CHO, RC(=O)R’; R≤4) produced from biomass

pyrolysis could be catalytically upgraded to value-added chemicals, hydrocarbons, or aromatics

[1-3] by reactions that remove oxygen heteroatoms and lengthen their carbon backbone. The

aldol-type condensation reactions couple alkanal (R-CHO) or alkanone (RC(=O)R’) reactants to

increase their carbon chain length and eject an oxygen atom as H2O without the use of external

H2 [2, 4, 5]. The condensation reactions may occur in acidic and basic medium and, in the

homogenous phase, mechanistic pathways and catalytic functions of acid and base have been

well established [6]. The acid-catalyzed C-C bond formation [6-8] occurs via an initial keto-enol

tautomerization of alkanal (or alkanone) to form the conjugate enol. The sequential nucleophilic

attack of the alpha carbon in enol to the protonated carbonyl group of alkanal (or alkanone)

creates an intermolecular C-C bond, thus lengthening the carbon backbone and forming a beta-

hydroxy alkanal (or alkanone), also known as an aldol. The base-catalyzed C-C bond formation

[6] involves the formation of a resonance-stabilized enolate and its sequential nucleophilic attack

to the carbonyl group of an alkanal or alkanone to evolve the aldol. Both acid and base catalyzed

reactions share a common sequential dehydration step that transforms the beta-hydroxy alkanal

(or alkanone) to an alkenal or alkenone, respectively, to complete a catalytic turnover [6].

Similar reactions have been reported on acid sites [9-11], basic sites [12, 13], or bifunctional

acid-base site-pairs [14-17] immobilized within solid structures. Condensation reactions on solid

basic sites (e.g. on Mg-Al mixed oxide) form almost exclusively the expected primary

condensation products at low temperatures (353-413 K) [12, 13]. Mg-Al mixed oxides catalyze

heptanal and benzaldehyde reactions to form a mixture of cross condensation (jasminaldehyde)

and self condensation (2-n-pentyl-2-nonenal) products at selectivities of 67-80 % and 20-33 %,

respectively, at 403 K [12, 13]. Alkali ion-exchanged zeolites (Na-X, K-X, and Cs-X), alkali

treated alumina (KOH-Al2O3), and hydrotalcite [[Mg0.6Al0.4(OH)2](CO3)0.20·0.84H2O] convert

propanal to 2-methyl-2-pentenal and 3-hydroxy-2-methylpentanal (373 K) [12]. Similarly,

11

alkaline earth metal oxides (e.g. MgO and SrO) promote butanal condensation to form

predominantly the self condensation product (2-ethyl-2-hexenal) at carbon selectivities above 90 %

with a small amount of 2-ethyl-2-hexenol, heptanone, and 2-ethylhexanal at 573 K [14]. As the

temperature increases, ketonization, reverse α-addition, cracking, and decarboxylation reactions

begin to occur on basic sites (e.g. on MgO/SiO2 [14], SrO/SiO2 [14], and CexZr1-xO2 [18]) at

detectable rates relative to those of primary condensation reactions, as reported for propanal

reactions on MgO/SiO2 and SrO/SiO2 (> 723 K) [14] and butanal reactions on CexZr1-xO2 (> 673

K) [18].

On Brønsted acid catalysts (H-MFI) [4, 5, 19], the initial coupling of alkanals leads to larger

oxygenates, which undergo secondary reactions of aromatization, alkylation/dealkylation, and

cracking. At higher temperatures (e.g. 673 K), these secondary reactions occur much faster than

the initial alkanal turnovers and lead to diverse methyl- or ethyl-substituted aromatics (e.g.

trimethylbenzene, methyl-ethylbenzene) and light gases (e.g. CO, CO2, C1-C3 hydrocarbons).

The primary condensation products were detected on H-MFI at these higher temperatures only

when introducing alkanal reactant in pulses to maintain the pressures of alkanal and primary

products at low values [4].

Condensation, aromatization, and cracking reactions during alkanal conversions on Brønsted

acid sites (H+) contained within microporous crystalline silica-alumina frameworks have been

recently proposed to involve pools of oxygenate and hydrocarbon intermediates co-existed

within the zeolitic pores [4]. The primary and secondary nature of these reactions, the rate

dependencies of individual catalytic paths, the identity of kinetically relevant steps, and the

associated mechanistic details have not been rigorously established. The lack of such molecular

level details is caused, in large part, by the complexity of the reaction network. The mechanistic

knowledge, kinetic dependencies, and site requirements for the C-C bond formation and oxygen

removal are, however, crucial for predicting and tailoring the product distributions and their

yields.

Herein, I interrogate the fate of propanal during its catalytic sojourns over Brønsted acid sites

(H+) contained within MFI framework. I propose a sequence of elementary reactions to describe

the fate of propanal in competitive reactions that create an inter- and intramolecular C=C bond to

evolve 2-methyl-2-pentenal and propene, respectively. I also discuss secondary reactions of ring

12

closure, dehydration, and dehydrogenation that evolve the larger aromatics and requirements of

hydrogen transfer between the secondary dehydrogenation step and the intramolecular C=C bond

formation step to satisfy the reaction stoichiometry. I draw a mechanistic synergy between the

reactions occurring at acid sites contained within the microporous crystalline materials and those

in the homogeneous phase and then report a competitive direct deoxygenation path undetected in

homogeneous reactions. The rate dependencies for these reactions were measured, their kinetic

relevance and reversibility were interrogated, and the kinetic couplings of hydrogen transfer

within the catalytic sequence were confirmed.

2.2. Experimental

2.2.1. Catalyst preparation

MFI zeolite samples in their NH4+ form (425 m

2·g

-1, Si/Al atomic ratio=11.5, CBV2314, Zeolyst)

were treated in flowing dry air (0.6 cm3·g

-1∙s

-1, zero grade, Linde) to 873 K by increasing the

temperature at 0.0167 K∙s-1

and holding for 4 h to convert NH4+ to H

+. In a separate series of

MFI zeolite samples, ion exchange was performed to exchange the NH4+ ions with Na

+ ions to

attain final samples with varying H+ and Na

+ contents (atomic ratios of H

+/Al=0.82-0.48 and

Na+/Al=0-0.34). NH4

+-MFI zeolites (Si/Al=11.5, 4 g) were mixed and stirred with aqueous NaCl

solution [7.8-70 g of NaCl (99 %, ACP Chemicals) and 100 cm3 of doubly-deionized water] at

298 K for 24 h. The samples were then washed in doubly deionized water and filtered under

vacuum until Cl─ ions in the filtrate were undetected from chemical titration with 0.1 mol·L

─1

AgNO3 (99.9999 %, Sigma Aldrich) aqueous solution. The samples were dried at 393 K for at

least 8 h and then heated in flowing dry air (0.6 cm3∙g

─1∙s

─1, zero grade, Linde) at 0.0167 K∙s

−1 to

873 K and holding for 4 h at 873 K to convert NH4+ to H

+.

2.2.2 Catalytic rates and selectivities of propanal and 1-propanol reactions on MFI zeolites

Chemical turnover rates and selectivities of propanal reactions were measured in a fixed bed

tubular microcatalytic quartz reactor (inner diameter of 9.5 mm) with plug-flow hydrodynamics

and operated under continuous flow mode. The reactor was contained within a resistively heated

furnace with its temperature regulated using a digital feedback temperature controller. Catalyst

13

samples (300 mg) were supported on a quartz frit and the reaction temperature was recorded

using a K-type thermocouple placed at the center of the axial and radial directions of the catalyst

bed.

Catalysts were heated in flowing He (2.8 cm3∙g

-1∙s

-1, Grade 5.0, Linde) at 0.0167 K∙s

-1 to reaction

temperatures (473-673 K) prior to rate measurements. Propanal (Kosher grade, ≥97 %, Sigma

Aldrich) was used as received and introduced via a gas tight syringe (Model 008025, 1 cm3, SGE)

mounted on a syringe infusion pump (Model LEGATO 100, KD Scientific) into a vaporizing

zone, in which it was evaporated and mixed with a He purge stream (0.83 cm3·s

-1, Grade 5.0,

Linde) at 320 K. All gas lines for transferring reactant mixtures were heated to 320 K and

product mixtures were heated to 473 K to prevent condensation. Reactions of 1-propanol

(≥99.9 %, Sigma Aldrich) on H-MFI were carried out following the same procedure with 0.58

kPa 1-propanol and at a space velocity of 6×10−4

mol 1-propanol·(mol H+

i·s)−1

.

Reactions using propanal and water (C3H6O-H2O) or propanal and 3-methyl-1-pentene (C3H6O-

C6H12) feed mixtures were carried out by introducing either doubly-deionized H2O or 3-methyl-

1-pentene (99 %, Sigma Aldrich) into a second vaporizing zone maintained at 363 K or 330 K,

respectively, located downstream from the vaporizer used for propanal evaporation (described

above). Doubly-deionized H2O or 3-methyl-1-pentene was introduced via a gas tight syringe

[Model 1005, 5 cm3 (Hamilton) for H2O and Model 006230, 0.25 cm

3 (SGE) for C6H12 infusion]

mounted on a syringe infusion pump (Model KDS-100, KD Scientific, for H2O or Model

LEGATO 100, KD Scientific, for C6H12 infusion).

Chemical species in the feed and reactor effluent stream were quantified using an on-line gas

chromatograph (Model 7890A, Agilent) and mass spectrometer (Model 5975C, Agilent), GC-

MS, equipped with a 10-port sampling valve containing two sample loops of 250 μl each. The

samples contained in the gas sampling loops were analyzed by chromatographic separation using

two capillary columns (HP-5MS, Agilent, 190091S-433, 30 m, 0.25 mm ID and HP-5, Agilent,

19091J-413, 30 m, 0.32 mm ID). The HP-5 column is connected to thermal conductivity (TCD)

and flame ionization (FID) detectors in series and the HP-5MS column to the MS detector.

14

2.2.3. Chemical titration of Brønsted acid sites

The number of Brønsted acid sites present initially (denoted by H+

i; herein and after, subscript “i”

represents the initial H+ site density) and remaining after reactions (denoted by H

+r, refer to

working H+ sites without binding to larger, inactive carbonaceous species) was determined from

pyridine titration at 473 K on samples treated in flowing dry air (Sec. 2.2.1) or after rate

measurements (Sec. 2.2.2), respectively. Acid site titrations on air treated MFI samples were

carried out on 300 mg samples after in-situ heat treatment to 473 K at a constant heating rate of

0.0167 K·s−1

under flowing He (0.83 cm3∙s

−1). Acid site titrations on samples after propanal

reactions were carried out upon the removal of propanal feed and purging under flowing He

(0.83 cm3∙s

−1) at 473 K for 1800 s. After these respective treatments, pyridine was introduced at

3.42×10−8

mol·s−1

via a gas tight syringe (Model 006230, 0.25 cm3, SGE) into a vaporization

zone maintained at 391 K, in which it was evaporated and mixed with a flowing He stream. The

pyridine/He mixture was then introduced to the sample and the amount of pyridine in the effluent

stream was quantified using the flame ionization (FID) detector on the gas chromatograph

(Model 7890A, Agilent). The titration was completed when the molar flow rate of pyridine in the

effluent stream became identical to that of the feed stream. The number of H+ site was

determined based on the pyridine uptakes assuming a pyridine-to-H+ molar ratio of unity. The

relative amount of Brønsted and Lewis acid sites was also determined from integrated intensities

of pyridine adsorption bands from infrared spectrum of pyridine adsorption on the H-MFI

samples at 1540 cm-1

and 1450 cm-1

, respectively, at 423 K.

2.2.4. Temperature programmed desorption of surface intermediates after propanal reactions on MFI zeolites

Temperature programmed desorption (TPD) was performed on the catalyst (300 mg MFI zeolite)

after propanal reactions [1.1 kPa C3H6O, 1.1×10−3

mol C3H6O·(mol H+

i·s)−1

] for 960 s at 473 K

in the microcatalytic quartz reactor described in Section 2.2.2, followed by purging under

flowing He (0.83 cm3∙s

−1) for 1800 s at 473 K to remove residual propanal from the transfer lines.

The sample was heated under flowing He (0.83 cm3∙s

−1) at a constant heating rate of 0.0167

K∙s−1

to increase the catalyst bed temperature from 473 to 673 K linearly and held isothermally

at 673 K for 1800 s. The composition of the effluent stream was analyzed using the GC-MS

(described in Sec. 2.2.2).

15

2.3. Results and discussion

2.3.1. Reaction network and product distributions during catalytic deoxygenation of propanal on H-MFI zeolites

Figures 2.1 and 2.2 show turnover rates and product distributions, respectively, for propanal

reactions on H-MFI zeolites (Si/Al=11.5) between 473 and 673 K and Appendix Section 2.6.1

summarizes the associated carbon balances. At 473 K, propanal reactions formed light alkenes

(predominantly propene), larger oxygenates (predominantly 2-methyl-2-pentenal and 2,3,4,5-

tetramethyl-2-cyclopentenal), aromatics (predominantly trimethylbenzenes and

dimethylbenzenes), and a small amount of CO and CO2 [ 2CO CO overall/r r r <0.01, 473 K, where

2COr , COr ,

overallr denote the CO2 and CO formation rates and the overall propanal conversion rate,

respectively]. As the reaction temperature increased from 473 to 548 K, the carbon fractions of

aromatic and olefinic species in the reactor effluent stream increased from 13.1 to 51.4 % and

from 9.5 to 33.4 %, respectively, while the carbon fraction of oxygenates concomitantly

decreased from 65.2 to 10.2 % (Figs. 2.2a-2.2c). At 673 K, oxygenates were undetected, because

of their rapid conversion to aromatic and olefinic species, which became the predominant

products with carbon distributions of 81.6 % and 18.3 %, respectively (Fig. 2.2d). Within the

olefinic and aromatic fractions, the species diversity increased with increasing temperature, and

the predominant species were propene, toluene, xylene, methylnaphthalene, and

dimethylnaphthalene as the temperature reached 673 K, as shown in Figure 2.2d.

16

Figure 2.1. Temperature dependence of propanal conversion rates (○) and the rates of olefin

(C2=-C6

=, ×), C6H10O (2-methyl-2-pentenal, ●), C9H14O (2,3,4,5-tetramethyl-2-cyclopentenal and

isomers, ▲), aromatic (C6-C12, ■) formation during propanal (C3H6O) reactions on H-MFI

zeolites [Si/Al=11.5, 7.5 ks, 1.1×10−3

mol C3H6O·(mol H+

i·s)−1

(subscript i denotes the initial

acid site density), 1.1 kPa C3H6O in He].

1.5 1.6 1.7 1.8 1.9 2.0 2.1

10-9

10-8

10-7

10-6

10-5

Total rate

Olefin

C6H10O

C9H14O

Aromatics

Re

actio

n r

ate

s (

mo

lg

ca

t.s

)1)

1000K/T

17

Figure 2.2. Carbon distributions in the effluent stream of propanal (C3H6O) reactions on H-MFI

zeolite (Si/Al=11.5) at 473 K (a), 523 K (b), 548 K (c), and 673 K (d) (7.5 ks, 1.1×10−3

mol

C3H6O·(mol H+

i·s)−1

, 1.1 kPa C3H6O in He, overallr is the overall C3H6O conversion rate).

18

Propanal conversion rates increased with temperature to values much larger than expected from

extrapolation of the rate data at lower temperatures (473-548 K, Fig. 2.1) using the Arrhenius

relation. Rates at 673 K exceeded the expected values by at least two orders of magnitude. These

temperature effects on rates indicate that additional reaction pathways, apparently those with

higher effective activation energies, become the predominant propanal conversion routes at

higher temperatures. The higher temperatures and prevalent higher pressures of aromatics

(mono-, di-, and trimethylbenzenes) and alkenes (ethylene, propylene, and butene) promote

additional reaction pathways, which may include the hydrocarbon pool typed mechanism, as

established previously for methanol-to-olefin or methanol-to-gasoline synthesis on MFI zeolites

[20-23].

Scheme 2.1 shows a proposed reaction network for propanal reactions on Brønsted acid sites.

This reaction network captures the competitive bimolecular and unimolecular reaction pathways

that create the inter- and intramolecular C=C bonds to evolve 2-methyl-2-pentenal (C6H10O) and

propene (C3H6), respectively, and the sequential ring closure and alkylation-dealkylation

reactions that form larger aromatics and oxygenates (> 6 carbon atoms). Bimolecular reactions of

propanal create intermolecular C-C bond before H2O elimination to form 2-methyl-2-pentenal as

the primary product (R 1.1, Scheme 2.1). The 2-methyl-2-pentenal undergoes sequential

coupling with propenol (from propanal tautomerization step, R 2.1) and H2O elimination (R 1.2)

to evolve 2,4-dimethyl-2,4-heptadienal (C9H14O). The 2,4-dimethyl-2,4-heptadienal may

undergo a ring closure step followed by isomerization (R 4.1) to form predominantly 2,3,4,5-

tetramethyl-2-cyclopentenal and a small amount of oxygenated isomers or, alternatively, a ring

closure and H2O elimination step (R 1.2.1) to create an intramolecular C=C bond and evolve

1,3,5-trimethylbenzene (C9H12). Sequential alkylation and dealkylation of trimethylbenzene (R 3)

lead to diverse aromatic species containing 6-12 carbon atoms (carbon distributions shown in Fig.

2.2). In parallel to the bimolecular reaction (R 1.1), propanal may undergo deoxygenation

reaction to remove the oxygen heteroatom by its recombination with two external H atoms to

eject H2O, forming an intramolecular C=C bond to evolve propene (R 2.1 and R 2.2). At higher

temperatures, secondary reactions of dehydrogenation, hydrogen transfer, and alkylation-

dealkylation steps that shuffle methyl groups and hydrogen atoms become much faster than the

initial inter- and intramolecular C=C bond formation in propanal (R 1.1 and R 2.1-2.2) and

19

secondary condensation reactions (R 1.2-1.3). These secondary reactions lead to diverse olefinic

and aromatic products with carbon distributions shown in Figure 2.2. In the next sections, I

provide the mechanistic evidence on the primary and secondary nature of the reaction steps

proposed in Scheme 2.1. I first probe and then confirm the accessibility of all Brønsted acid sites

that are contained within the MFI framework to propanal reactants in Section 2.3.2 and then

interrogate the specific catalytic requirements and kinetic dependencies for the competitive

bimolecular (R 1.1) and unimolecular (R 2.2) reactions that create the inter- and intramolecular

C=C bond within propanal to evolve 2-methyl-2-pentenal and propene, respectively, in Sections

2.3.3 and 2.3.4. The reversibility of the kinetically relevant steps and the kinetic relevance of

hydrogen transfer are discussed in Sections 2.3.5 and 2.3.6.

Scheme 2.1. Reaction network for propanal turnover on H-MFI zeolite (#: Intermediates at

undetectable concentrations; *: Hydrogen donating agents for deoxygenation reaction in R 2.2).

R 2.1 R 2.2

R 1.1

R 1.2

R 1.3

R 1.2.1

R 1.3.1 R 1.3.2 R 1.3.3

R 3R 4.1

R 4.2

PropanalPropenol

Propene

Propenol

2,4-dimethyl-2,4-

heptadienal

2,3,4,5-tetramethyl-

2-cyclopentenal

2- methyl-2-pentenal

Propenol

1,3,5-trimethylbenzene

Propenol

2,4,6-trimethyl-

2,4,6-nonatrienal

2,6-di-isoproylphenol

Alkylbenzenes

Olefins

6,8-dimethyl-

1,2,3,4-tetrahydronaphthalene

#

# **

6,8-dimethyl-

1,2-dihydronaphthalene

1,3-dimethyl-

naphthalene

(from H donating agents, e.g. 6,8-dimethyl-

1,2,3,4-tetrahydronaphthalene,

Steps R 1.3.2 and R 1.3.3)

Propene

(from R 2.2)

Propanal

Propanal

R 2.1#

#

##

PNG

20

2.3.2. Accessibilities of Brønsted acid site to propanal reactant and effects of acid site density on propanal conversion rates

Infrared spectroscopic studies of pyridine adsorption on the H-MFI (Si/Al=11.5) samples used in

this study shows pyridine adsorption on Brønsted acid sites at 1540 cm-1

and Lewis acid sites at

1450 cm-1

. The integrated intensities of these peaks, after correction with their respective

extinction coefficients, show that 86.7 % of the total sites are Brønsted acid sites while the rest

are Lewis acid sites. H2O byproducts from inter- and intramolecular C=C bond formation

reactions either rehydrate the unsaturated aluminum sites (Lewis sites) on non-framework

alumina to form surface hydroxyl species or physisorb on these sites [9, 24], thus preventing

them from participation in catalysis. Thus, steady-state propanal conversion rates (> 7.5 ks)

reflect predominantly the contributions from Brønsted acid sites, their amount available for

propanal catalysis was quantified by pyridine titration after exposure of the catalysts to reaction

mixtures.

The vast majority of Brønsted acid sites (H+) participating in catalytic turnovers are

predominantly occupied by propanal (C3H6O) and its isomers during catalysis at 473 K and 1.1-

4.5 kPa of C3H6O, a condition required for the rates of inter- and intramolecular C=C bond

formation to vary proportionally to and remain independent of propanal pressure, respectively, as

shown later in Sections 2.3.3 and 2.3.4. The requirements of propanal and its isomers as the most

abundant surface intermediates are independently confirmed next from temperature programmed

desorption carried out on H-MFI (Si/Al=11.5) after propanal catalysis at 473 K [1.1 kPa C3H6O,

960 s, 1.23×10−6


].

Figure 2.3 shows the temperature programmed desorption profile, plotted as the rate of

desorption of carbonaceous species (per H+

i site) from the H-MFI sample versus temperature.

Aromatics (94 % C8~C12) and a small amount of light alkenes (6 % CnH2n, n=2-3) are the

predominant species desorbed from the sample, because adsorbed intermediates would undergo

intermolecular C=C bond formation and ring closure steps (Pathways R 1.2-1.2.1 and R 1.2, 1.3,

and 1.3.1, Scheme 2.1) before their desorption as larger aromatic species or cracking of these

larger species to form the small amount of alkenes. Integration of the desorption rates in Figure

2.3 over the entire temperature range gives the cumulative amount of carbon desorbed from the

21

H-MFI catalyst. This amount, in terms of carbon-to-H+ (C/H

+) atomic ratio, was found to be 3.15.

The value translates to a C3H6O-to-H+ molar ratio of near unity (1.05) during steady-state

catalysis. This value, taken together with the first-order rate dependence for the intermolecular

C=C bond formation (Sec. 2.3.3) and zero-order dependence for the intramolecular C=C bond

formation (Sec. 2.3.4) in propanal, is consistent with binding of propanal or its isomers to H+

sites as the most abundant surface intermediates. The value infers that all H+ sites were

accessible to propanal.

Figure 2.3. Desorption rate of carbonaceous species from H-MFI catalyst as a function of

temperature. The temperature programmed desorption was performed after exposure of the

catalyst to propanal (C3H6O) reactions for 960 s at 473 K (300 mg H-MFI, Si/Al=11.5, 0.0167

K∙s−1, propanal reaction conditions: 1.1 kPa C3H6O in He, 1.23×10

−6 mol C3H6O·(gcat.·s)

−1).

All of the H+ sites in the 10-membered ring MFI channels (diameter of ~0.55 nm) are accessible

to propanal, a compound with an estimated kinetic diameter of 0.45-0.50 nm (based upon the

kinetic diameters for alkane/alkene/alcohol with three or four carbon atoms [25, 26]), because

these sites are accessible to molecules with comparable (C3-C6 n-alkanes, 0.43-0.50 nm [27])

and larger (pyridine, used here as the titrant, 0.59 nm [28]) kinetic diameter values. The

accessibilities of all H+ sites to propanal are also confirmed from rate measurements on a series

of MFI zeolites (Si/Al=11.5) with different H+-to-Na

+ ratios (molar ratios of H

+/Al and Na

+/Al

500 550 600 650 700 7500

2

4

6

8

10

Temperature (K)

Constant

673K

temperature

De

so

rptio

n r

ate

of

ca

rbo

na

ce

ou

s s

pe

cie

s

(104m

ol C(

mo

l H

+ is

)1)

22

are 0.82-0.48 and 0-0.34, respectively). The overall rates of propene formation (rintra) from the

primary deoxygenation route (R 2.1 and 2.2, Scheme 2.1) are plotted as a function of H+ and Na

+

site densities in Figure 2.4. The rates of propene formation (per mass of MFI sample) decreased

proportionally with the decrease in H+ site density and thus the increase in Na

+ site density, as

these site densities were interrelated (the sum of H+ and Na

+ site density equals 82 % of the

nominal Al density). Extrapolation of the propene formation rates to zero H+ site density gave

propene formation rates below detectable values because substitution of H+ with Na

+ ions

removed the Brønsted acid sites required for the direct deoxygenation reaction (R 2.2, Scheme

2.1). Rates for propene formation [per mass of MFI sample, rintra (per gcat.)] are strictly

proportional to H+ site densities with an extrapolated rate value of zero for MFI samples without

any H+ site (Fig. 2.4). These rate values (per mass of MFI sample) led to constant turnover rates

(per H+ site) of 1.42×10

−6±0.03×10

−6 mol∙(mol H

+i∙s)

−1 (473 K) for propene formation,

irrespective of the initial H+ site density (0.72×10

−3−1.12×10

−3 mol H

+i∙gcat.

−1). Substitution of H

+

with Na+, however, increased the rates of bimolecular reaction (R 1.1, Scheme 2.1), because Na

+

sites catalyze propanal condensation much more effectively than H+ sites via a separate, base

catalyzed aldol condensation pathway. Framework oxygen ions conjugated to Na+ sites act as

weak bases [29] and abstract the α-hydrogen on propanal to form carbanion intermediates [30],

which undergo nucleophilic addition to the carbonyl group of another propanal and, upon

dehydration, create the intermolecular C=C bond. The constant turnover rate values for propene

formation irrespective of H+ site density (Fig. 2.4), together with a C3H6O-to-H

+ surface molar

ratio of 1.05 after catalysis (Fig. 2.3), led us to conclude that all H+ sites within the MFI zeolites

were accessible to propanal and these sites were kinetically equivalent (for propene turnovers).

Next, I report kinetic dependencies and catalytic requirements for the inter- and intramolecular

C=C bond formation, interrogated under conditions that minimize both the secondary reactions

(alkylation, cracking, etc.) and involvement of hydrocarbon pool mechanism. These conditions

were attained by rate measurements under differential conditions (<5 % propanal conversion), a

high space velocity [1.1×10−3

mol C3H6O·(mol H+

i·s)−1

], and at a moderate temperature (473 K).

23

Figure 2.4. Overall rates (per mass of catalyst, ●) and turnover rates (per H+

i, subscript i denotes

the initial acid site density, ○) for intramolecular C=C bond formation in propanal (C3H6O) on H-

MFI plotted as a function of H+

i and Na+ concentration (473 K, Si/Al=11.5, 7.5 ks, 1.23×10

−6


, 1.1 kPa C3H6O in He).

2.3.3. Kinetic dependencies, elementary steps, and site requirements for intermolecular C=C bond formation of propanal on H-MFI zeolites

Rates for bimolecular propanal reaction that creates an intermolecular C=C bond and lengthens

the carbon chain (R 1.1, Scheme 2.1) were measured at different propanal pressures (1.1-4.5 kPa)

on H-MFI (Si/Al=11.5) at 473 K. Turnover rates for this reaction, inter

r , were calculated based

on the number of H+ site available during reaction (determined from pyridine titration right after

the rate measurement, Sec. 2.2.3), because condensation reactions of propanal produce heavy

compounds (e.g. naphthalenes) which can gradually deposit on the catalysts and reduce the

amount of H+ sites available for catalytic turnovers, especially at the longer times (e.g. 7.5 ks for

rate data reported herein). The rate values at 7.5 ks, which reflect the steady-state catalytic rates,

are plotted as a function of C3H6O pressure in Figure 2.5. Turnover rates for intermolecular C=C

0 200 400 600 800 1000 12000.0

0.5

1.0

1.5

Re

actio

n r

ate

(1

0-8m

ol(

gca

t.s

1)

H+

i-site density (10

6molg

cat.

1)

1000 800 600 400 200 0

0

1

2

3

4

rintra

(per H+

i)

rintra

(per gcat.

)20

Re

actio

n r

ate

(1

06m

ol(

mo

l H

+ is

)1)

Na+-site density (10

6molg

cat.

1)

24

bond formation (inter

r ) increased linearly with C3H6O pressure (3 6C H OP =1.1-4.5 kPa), following

the expression of:

3 6C H Ointer inter, effPr k (2.1)

where kinter,eff is the effective rate constant for intermolecular C=C bond formation. I next

propose a sequence of elementary steps, as presented in Scheme 2.2, from which I derive a rate

expression that accurately describes the observed first-order rate dependence on C3H6O pressure.

Figure 2.5. Turnover rates for intermolecular C=C bond formation (rinter, ■) and intramolecular

C=C bond formation (rintra, ●) in propanal (C3H6O) that evolve 2-methyl-2-pentenal (C6H10O)

and propylene, respective, and the rate ratio for inter- over intramolecular C=C bond formation

(rinter/rintra, ○), plotted as a function of C3H6O pressure on H-MFI at 473 K [Si/Al=11.5, 7.5 ks,

1.1×10−3

-4.4×10−3

(mol C3H6O·(mol H+

i· s)−1

)]

0 1 2 3 4 50

2

4

6

8

rinter/rintra

rinte

r/rin

tra

Tu

rno

ve

r ra

te (

10

-4m

ol

·(m

ol H

+ r·s

)-1)

PC3H6O

(kPa)

0

5

10

15

20

25

30

rintra

10

rinter

25

Scheme 2.2. Mechanism for inter- and intramolecular C=C bond formation in propanal (C3H6O)

evolving 2-methyl-2-pentenal (C6H10O) and propylene, respectively, on H+ sites ( denotes

quasi-equilibrated step, reversible step, and irreversible step). In taut,' sK of Step

2.1b, species s denotes propanal surface isomers [(4a), (4b), and their physisorbed isomers].

Propanal (1) and propenol (2) interconvert rapidly within the intra-zeolitic channels (Step 2.1a,

Scheme 2.2), thus these species are treated as a kinetically indistinguishable chemical lump and

their relative pressures are dictated by the thermodynamics of keto-enol tautomerization reaction.

Propanal may adsorb at the Brønsted acid sites (H+) in diverse isomeric configurations of

protonated propanal (3), propenol (4a), allyl alcohol (4b), or in the physisorbed forms of these

species via single or multiple hydrogen bonds between the H+ and the oxygen atom in these

species and/or between the lattice oxygen and the hydrogen on propanal isomers. The adsorption

configurations of similar carbonyl compounds (e.g. acetone) on acidic zeolites (MFI [31-35], Y

[32]) have been probed using 13

C NMR studies. The formation of surface enol has been

previously confirmed from H/D exchange between adsorbed acetone-d6 and H+ sites and

between 13

C-2-acetone and D+ sites in MFI zeolites at ambient temperature [31] and from the

(2)

Propenol (g)

(3)

Propanal (ads)

Step 1.1

Step 1.2

Step 1.3 Step 1.4

Step 2.2

Ste

p 2

.1a

Step 2.3 Step 2.4

Ste

p 2

.1b

(1)

Propanal (g)

(4a)

Propenol(ads)

(5)

3-hydroxyl-2-

methylpentanal (ads)

(4b)

Allyl alcohol (ads)

(6)

2-methyl-

2-pentenal (ads)

(8)

Propanol (ads)

(7)

2-methyl-2-pentenal (g)

(10)

Propylene (g)

Step 2.5

(9)

Propoxide (ads)

,K C H O6 10ads

K, C H3 6ads

Ste

p 2

.1a

(1)

Propanal (g)

(from H donating agents)

or

other physisorbed propanal isomers

(2)

Kdehy

KalkoxkH trans

kaldol

kaldolKads

'K ads

' sK taut, Ktaut

Ktaut

26

appearance of signal at ~180 ppm on 13

C NMR spectra as acetone adsorbed on H-MFI and H-Y

between 298 and 453 K [32]. The formation of hydrogen bonded propanal from the less stable

allyl alcohol was confirmed from the feature at 216 ppm in 13

C NMR spectrum emerged during

adsorption of allyl alcohol on H-MFI between 400 and 425 K [33]. Density functional theory

(DFT) calculations on alcohol (CnH2n+1OH, n=1-4) adsorption in H-MFI structures show a broad

and shallow potential energy well between the protonated and hydrogen bonded species with

small energy differences (between −1 and +6 kJ∙mol−1

for CnH2n+1OH, where n=1 [36-38], 2 [38],

3 [38], and 4 [38, 39]). Similar proton affinities between alcohol and alkanal (e.g. 786.5 kJ·mol─1

and 786 kJ·mol─1

for n-propanol and propanal, respectively [40]) suggest that the proton transfer

step may also occur readily in alkanals, thus species (3), (4a), (4b), and their physisorbed

counterparts may interconvert rapidly. Equilibrium between propanal and propenol (Step 2.1a)

and between these gas phase species and their respective adsorbed complexes (Steps 1.1 and 2.2)

requires complexes (3), (4a), (4b), and other adsorbed propanal isomers to remain equilibrated

with each other within the time-scale of forward propanal turnovers. Rapid interconversion

among the various propanal derived surface intermediates renders these species be treated as a

kinetically indistinguishable lump, and within this lump, their relative surface abundances are

dictated by the differences in their heats of adsorption.

Nucleophilic addition of propenol (2) to the carbonyl group of protonated propanal (3) (Step 1.2)

creates an intermolecular C-C bond and produces a 3-hydroxy-2-methylpentanal (5). Sequential

H2O elimination (Step 1.3) converts the newly formed C-C bond in the 3-hydroxy-2-

methylpentanal (5) to a C=C bond and increases its degree of unsaturation, and upon desorption

(Step 1.4), completes a catalytic turnover to evolve 2-methyl-2-pentenal (7). These mechanistic

pathways are analogous to those established in the liquid phase [6, 7] and proposed for

condensation reactions of acetone on H-MFI samples [41]. The nucleophilic addition step (Step

1.2) is the kinetically relevant step and this step is reversible, as shown in Section 2.3.5. Quasi-

equilibrium assumptions for Steps 1.1, 1.3, 1.4, 2.1a, 2.1b, 2.2, 2.4, and 2.5 in Scheme 2.2,

together with pseudo-steady-state approximation on all reactive intermediates, lead to the

following rate expression for intermolecular C=C bond formation (rinter, per catalytically active

H+ site):

27

2

2

6 10

3 6 6 10

6 10 3 6

3 6 3 6 6 10 6 10 3 6 2 3 66 10 3 6

ads,C H O2taut C H O C H O H Oaldol ads aldol

dehy

interads,C H O ads,C H

C H O C H O C H O H O C H O C H H O C Hads ads ads,C H O ads,C H1 dehy alkox

taut,1 '

n

ss

Kk K K P k P P

Kr

K KK P K K P P P K P P P K P

K K

+ + + + + +5 7 73 6 3 6 12 2 6 10 3 3 (H ) (C H O-H ) (C H OH-H ) (C H O -H ) (C H O-H ) (C H OH-H ) (C H -O)

( ) ( ), ( ), and their

3 4a 4b ( ) ( ) ( ) ( )

physisorbed forms

5 6 8 9

(2.2)

where rate and equilibrium constants and the associated catalytic steps are defined in Scheme 2.2.

kaldol and k−aldol denote the elementary rate constants for the forward and reverse reactions,

respectively, in Step 1.2. ads

K , tautK , taut,' sK , dehy

K , alkox

K , 6 10C H Oads,

K , and 3 6C Hads,

K

denote the equilibrium constants for Steps 1.1, 2.1a, 2.1b, 1.3, 2.4, 1.4, and 2.5, respectively. xP

denotes the partial pressure of species x (x = C3H6O, C6H10O, C3H6, or H2O). The diverse

propanal surface isomers [(4a), (4b), and their physisorbed forms] interconvert rapidly with each

other; thus, their overall surface concentration over that of vacant H+ sites is given by term

3 6C H Otaut,ads1

'n

ss

K K P over unity, where taut,' sK is the equilibrium constant for the formation of

species s [s represents (4a), (4b), and the physisorbed forms of (4a), (4b), etc.] from protonated

propanal (3) (Step 2.1b, Scheme 2.2), thus term 1

taut,'n

ssK

represents the aggregated

equilibrium constant for the adsorbed propanal isomers. The magnitude of each term in the

denominator of Equation 2.2 reflects the surface coverage ratio of a specific surface species [(3),

((4a), (4b) and their physisorbed forms), (5), (6), (8), and (9), as labeled in the equation] to

unoccupied H+ sites (the first term in the denominator with a value of unity). The rate expression

(Eqn. 2.2) is simplified to:

3 6 3 6C H O C H Otaut

inter aldol inter, eff

1taut,

1 'n

ss

Kr k P k P

K

(2.3)

when H+ sites are occupied by protonated propanal (3), propenol (4a), allyl alcohol (4b), and

their physisorbed forms as the most abundant surface intermediates and the reverse rates for the

intermolecular C=C bond formation, which are given by term

28

26 106 10

1C H O H Oaldol ads,C H O dehyk K K P P

in Equation 2.2, are insignificant compared with the net rates

of this reaction. Equation 2.3 accurately describes the first-order dependence measured and

presented in Figure 2.5, attained under differential conditions and thus at low C6H10O (7) and

H2O pressures (6 10C H OP =0.006~0.022 kPa,

2H OP =0.01~0.03 kPa). The effective rate constant in

Equation 2.1, inter,effk , equals the proportionality constant

1

taut taut,aldol1

1 'n

ss

k K K

in

Equation 2.3 and is the product of the forward rate constant for the nucleophilic addition step

( aldolk , Step 1.2), equilibrium constants for conversion between propanal and propenol ( tautK ,

Step 2.1a) and between protonated propanal and the various adsorbed propanal isomers ( taut,' sK ,

Step 2.1b). Equation 2.2 is used to regress against the rate data to extract the kinetic and

thermodynamic parameters, to be discussed in Section 2.3.7.

2.3.4. Kinetic dependencies, elementary steps, and site requirements for intramolecular C=C bond formation in propanal on H-MFI zeolites

Next, I provide the rate dependencies and mechanistic evidence for intramolecular C=C bond

formation in propanal by direct removal of oxygen heteroatom while preserving the carbon

backbone to evolve propene. The primary nature of this step is confirmed from the near

exclusive formation of alkenes (CnH2n, n=3-5) from alkanals (CnH2nO, n=3-5) following the

reaction stoichiometry of:

CnH2nO+2H CnH2n+H2O (2.4)

Within the lump of alkene product, the fraction of alkene with carbon number identical to the

alkanal reactant from C3H6O, C4H8O, and C5H10O reactions on H-MFI zeolites, i.e.

[C3H6]/[CnH2n], [C4H8]/[CnH2n], and [C5H10]/[CnH2n], are 0.93, 0.95, and 0.89, respectively, at

473 K and high space velocity [1.1×10−3

mol CnH2nO·(mol H+

i·s)−1

] at which sequential

conversions of alkene remain negligible [42].

29

Turnover rates for intramolecular C=C bond formation (intra

r ) during propanal reactions on H-

MFI (Si/Al=11.5) at 473 K are essentially insensitive to propanal pressure (1.1-4.5 kPa), as

shown in Figure 2.5, and equal to the effective rate constant intra,effk :

0

3 6C H Ointra intra, effr k P (2.5)

This rate dependence is consistent with the proposed mechanism in Scheme 2.2. Propanal

molecules first encounter H+ sites and adsorb as protonated species (3), which interconvert to

diverse isomeric surface species [(4a), (4b), and their physisorbed isomers], as described in

Section 2.3.3. As these isomeric species interconvert, a portion of them accepts hydrogen from

hydrogen donating agents (identity and origin to be discussed in Sec. 2.3.6) in a kinetically

relevant hydrogen transfer step (Step 2.3) to increase their degree of saturation and form

adsorbed propanol (8). The adsorbed propanol undergoes sequential dehydration (Step 2.4) to

form surface propoxide (9) [43-46] before desorption as propene (10) (Step 2.5). The

dehydration step occurs after the hydrogen transfer step and remains kinetically inconsequential

(to be confirmed in Sec. 2.3.6). These assumptions, together with pseudo-steady-state

approximation applied to the various surface intermediates in Scheme 2.2, give the following

rate expression for intramolecular C=C bond formation (rintra, per catalytically active H+ site):

6 10 3 6

6 10 2 6 10 3 6 2 3 66 10 3 6

2 3 6

3 63 6

H trans R'Hads1

intraads,C H O ads,C H

C H O C H O H O C H O C H H O C Hads ads ads, C H O ads,C H1 dehy alkox

+3 6

taut, C H O

taut, C H O

'

1 '

(H ) (C H O-H )

n

s

n

s

s

s

K K k P P

rK K

K P K K P P P K P P P K PK K

+ + + + +5 7 73 6 12 2 6 10 3 3 (C H OH-H ) (C H O -H ) (C H O-H ) (C H OH-H ) (C H -O)

( ) ( ), ( ), and their 3 4a 4b ( ) ( ) ( ) ( )

physisorbed forms

5 6 8 9

(2.6)

The elementary rate and equilibrium constants in Equation 2.6 are defined in the previous section

and provided in Scheme 2.2. H transk represents the effective rate constant for hydrogen transfer

(Step 2.3) and 2R'HP denotes the aggregate partial pressure of hydrogen donors, defined here as a

pool of species formed from secondary reactions (R 1.3.1 and 1.3.2, Scheme 2.1) that donate

hydrogen atoms (Step 2.3) to increase their extent of unsaturation. The kinetic relevance of

hydrogen transfer (Step 2.3), kinetic irrelevance of propanal dehydration step (Step 2.4), and

30

identity of hydrogen donors are probed and confirmed in Section 2.3.6. The term 2H trans R'Hk P in

Equation 2.6 is defined as:

2 2H trans H trans, R'H R'H , 1

m

jj

jk P k P

(2.7)

where 2R'H , jP and

H trans, jk are the pressure and hydrogen transfer rate constant for each specific

hydrogen donor j (e.g. 6,8-dimethyl-1,2,3,4-tetrahydronaphthalene, Scheme 2.1), formed from

secondary ring closure reactions (R 1.3.1, Scheme 2.1). The rates for intramolecular C=C bond

formation (Fig. 2.5) were measured at similar conversions (1.2-1.5 %) while individual pressures

of hydrogen donors were maintained at similar values (the total pressure of aromatics varied

from 4.5×10−4

to 6.5×10−4

kPa), thus 2H trans R'Hk P in Equation 2.6 is treated as a constant value in

our kinetic analysis (Sec. 2.3.7). Turnover rates for intramolecular C=C bond formation (intra

r ,

Eqn. 2.6) become independent of propanal pressure and their values equal the effective rate

constant for intramolecular C=C bond formation (intra,eff

k ), when H+ sites are predominantly

occupied by protonated propanal (3) and its isomers [(4a), (4b), and their physisorbed forms]

during steady-state catalysis:

2 3 6 3 6

0 01H trans R'H C H O C H Ointra

1

taut,

intra,eff

taut,

'

1 '

n

sn

s

s

s

K

r k P P k P

K

(2.8)

Equation 2.8 describes the measured rates for intramolecular C=C bond formation in Figure 2.5,

which remain largely insensitive to propanal pressure (with an apparent reaction order with

respect to C3H6O of 0.09±0.02). The slight increase in rates with propanal pressure reflects the

increase in hydrogen transfer probabilities as the pressures of hydrogen donors (2R'HP , limited to

a portion of the aromatic species produced, see Sec. 2.3.6) increase slightly with increasing

propanal pressure (a 4-fold increase in C3H6O pressure leads to ~50 % variation in the partial

pressure of total aromatics during rate measurements).

31

2.3.5. Reversibility of the inter- and intramolecular C=C bond formation in propanal on H-MFI zeolites

The influence of H2O on the number of catalytically active sites and on the net rates of inter- and

intramolecular C=C bond formation was probed by propanal reactions in excess H2O (5-10 kPa

H2O, H2O/C3H6O=4.5-9). Pyridine titration carried out after steady-state reactions with C3H6O-

H2O mixtures was used to determine the number of active sites that were free of larger, inactive

carbonaceous species and thus participated in catalytic turnovers. Table 2.1 summarizes the

pyridine uptakes measured after steady-state reactions in C3H6O-H2O mixtures [7.5 ks, 473 K,

0.01-9 H2O/C3H6O ratios, 1.1×10−3

mol C3H6O·(mol H+

i·s)−1

]. Pyridine uptakes were higher

when introducing H2O together with C3H6O reactant; their values paralleled the increase in

H2O/C3H6O ratio, because H2O scavenges the carbonaceous species, prevents site occupation,

and thus retards catalyst deactivation. Similar effects of H2O were found during methanol-to-

olefins reactions on similar catalysts (H-MFI [47-49]) and SAPO-34 [50, 51] at higher

temperatures (573-723 K).

Table 2.1. Pyridine uptakes on H-MFI zeolites from pyridine titration carried out after steady-

state reactions in C3H6O-H2O mixturesa

H2O pressure (kPa) H2O/C3H6O ratio Pyridine uptake (10−6

mol·gcat. −1

) Pyridine/H+

i ratio

~0.011 ~0.01 40 0.036

5 4.5 60 0.054

10 9 59 0.053

aSi/Al=11.5, 473 K, 7.5 ks, 1.1 kPa C3H6O, 0.011-10 kPa of H2O, space velocity=1.1×10−3 mol C3H6O·(mol H+i·s)−1. H+

i denotes

the number of Brønsted acid sites present initially (measured from pyridine titration after treating the H-MFI samples under

flowing He).

Turnover rates for inter- and intramolecular C=C bond formation in C3H6O-H2O mixtures are

plotted against the H2O pressure in Figure 2.6. The rates for intermolecular C=C bond formation

decreased from 2.2×10-4

mol∙(mol H+

r ∙s)-1

to 1.5×10-4

mol∙(mol H+

r∙s)-1

(the subscript “r”

represents H+ sites available for propanal turnovers, determined from pyridine titration, Sec.

2.2.3) as the H2O pressure increased from ~0.011 to 10 kPa. These effects of H2O reflect an

32

increase in the reverse rate of the nucleophilic addition step (Step 1.2, Scheme 2.2) at high H2O

pressures, as the numerator term [ 6 10 26 10

1

C H O H Oaldol ads, C H O dehyk K K P P

] of Equation 2.2

increases to a magnitude comparable to the term 3 6

2taut C H Oaldol adsk K K P of the same equation. As a

result, the net rate for intermolecular C=C bond formation in Equation 2.2 decreases and acquires

the form of:

2 26 10 6 10 6 10

3 6 3 6

3 6 3 6

1ads, C H O dehy C H O H O C H O H Otaut

C H O C H Ointer aldol aldol inter,eff inter,effC H O C H O

ads1 1

taut, taut,1 ' 1 '

n n

s ss s

K K P P P PKr k P k k P k

P PK K K

(2.9)

In contrast, the rates of intramolecular C=C bond formation did not vary at detectable extents and

maintained at 4.2×10-5

± 0.1×10-5

mol∙(mol H+

r∙s)-1

(Fig. 2.6). The lack of H2O effects indicates

that H2O as a byproduct does not alter the net rate of intramolecular C=C bond formation, thus

the kinetically relevant hydrogen transfer step is irreversible. The lack of H2O effects was also

found during allyl alcohol (a propanal isomer) conversion to olefins (predominantly C3H6) on

NaHY zeolites at 523 K [52], which may occur via similar mechanistic steps. The lack of H2O

effects on the intramolecular C=C bond formation also precludes the mechanistic sequence

involving reversible H2O elimination from propenol to evolve surface allylic alkoxides before

hydrogen insertion, because this case would lead the rates for intramolecular C=C bond

formation to decrease with increasing H2O pressure. The rate of olefin formation from

dehydration of alcohol on Brønsted acid catalysts (e.g. n-butanol on H-MFI [53], sec-butyl

alcohol on H-MFI [54], and 2-propanol on H-MOR [55]) was found to be zero-order at low

alcohol pressures (0.6-2 kPa). The zero-order rate constant for n-butanol dehydration on H-MFI

was estimated to be 2×10─2

mol∙(mol H+∙s)

─1 at 473 K, determined from extrapolation using the

Arrhenius relation from 378-458 K [53]. Dehydration of 1-propanol at conditions similar to the

propanal rate measurements reported here (0.55 kPa 1-propanol, 6×10−4

mol 1-propanol·(mol

H+

i·s)−1

, 473 K) led to complete conversion to propene and to a zero-order rate constant larger

than 1.5×10─2

mol∙(mol H+

r∙s)-1

at 473 K. Taken together the rate constant values for 1-propanol

and n-butanol dehydration at comparable conditions, I conclude that 1-propanol dehydration rate

must be at least two orders of magnitude larger than the direct propanal deoxygenation rate to

33

propene [1.5×10─4

-2.2×10─4

mol∙(mol H+

r ∙s)─1

] at 473 K, thus the H2O elimination step cannot

be the kinetically relevant step for the intermolecular C=C bond formation in propanal.

Figure 2.6. Effects of water on turnover rates for intramolecular C=C bond formation (intra

r , ●)

and intermolecular C=C bond formation (inter

r , ▲) in propanal (C3H6O) during C3H6O reactions

on H-MFI at 473 K (Si/Al=11.5, 7.5 ks, 1.1×10−3

mol C3H6O·(mol H+

i·s)−1

, 1.1 kPa C3H6O in

He).

These results led us to conclude that catalytic roles of H2O involve protecting the active sites

from occupation by larger, inactive species and altering the net rates for intermolecular C=C

bond formation by promoting the reverse reaction. H2O, however, does not alter the net rates for

intramolecular C=C bond formation because of its irreversible nature. Next, I probe the kinetic

relevance of hydrogen transfer for intramolecular C=C bond formation in propanal, the kinetic

coupling of this step to the secondary dehydrogenation reactions, and identify the hydrogen

transfer agents participated in the dehydrogenation steps within the reaction network.

0 2 4 6 8 100

5

10

15

20

25

30

rintra

Turn

over

rate

(10

-5m

ol·

(mol H

+ r·s

)-1)

PH2O (kPa)

rinter

34

2.3.6. Kinetic relevance of hydrogen transfer and requirements of hydrogen for intramolecular C=C bond formation in propanal

The direct deoxygenation route (Eqn. 2.4) for intramolecular C=C bond formation in propanal

requires oxygen removal from propanal by combining the oxygen heteroatom with two external

hydrogen atoms (R 2.2 in Scheme 2.1 and Step 2.3 in Scheme 2.2) to eject H2O. The rate ratios

for the overall CO and CO2 formation to propene formation were lower than 0.1 at 473 K

[ 2CO CO overall/r r r <0.01], thus the oxygen in propanal does not eject as CO or CO2 but instead

as H2O, as also confirmed from the near exclusive formation of propene ([C3H6]/[CnH2n]=0.93,

473 K) within the alkene fractions (Sec. 2.3.4).

A detailed structural analysis of the diverse aromatic and olefinic species in the effluent stream

led us to propose that a portion of hydronaphthalenes (e.g. 6,8-dimethyl-1,2,3,4-

tetrahydronaphthalene) with 10-15 carbons are the hydrogen donors via dehydrogenation steps

that increase their aromaticity (R 1.3.2 and 1.3.3, Scheme 2.1). The amount of H made available

from the dehydrogenation steps was determined to be 9.7×10−6

mol H∙(gcat.∙s)−1

at 673 K

[reaction conditions of 1.1 kPa C3H6O and 1.1×10−3

mol C3H6O·(mol H+

i·s)−1

], quantified here

based on the amount of aromatics (dihydronaphthalenes and naphthalenes) formed (Fig. 2.2d) as

a result of H donation (reactions R 1.3.2 and 1.3.3, Scheme 2.1) during steady-state catalysis.

This amount is consistent with the hydrogen amount required for the formation of diverse

olefinic species (R 2.2 and R 3, Scheme 2.1), calculated to be 9.0×10−6

mol H∙(gcat.∙s)−1

, based on

the alkene distributions in the product of the same reaction (Fig. 2.2d, 673 K). These results

indicate that intramolecular C=C bond formation requires H transfer to satisfy the reaction

stoichiometry, and in the absence of an external hydrogen source, the hydrogen must come from

secondary dehydrogenation events. The kinetic coupling of intramolecular C=C bond formation

and dehydrogenation steps thus dictates the distributions of olefinic and aromatic species formed

from propanal reactions.

The kinetic relevance of the hydrogen transfer step was probed by measuring rates in the

presence of 3-methyl-1-pentene (CH2=CH2CH(CH3)C2H5, denoted hereinafter as C6H12) as an

effective hydrogen donor [56], because of its weak tertiary allylic C-H bond (323 kJ∙mol−1

[57]).

Figure 2.7 shows the effects of C6H12 pressure on the turnover rates for inter- and intramolecular

35

C=C bond formation (inter

r and intra

r , respectively) in C3H6O-C6H12 mixtures on H-MFI at 473

K.

Figure 2.7. Effects of 3-methyl-1-pentene (C6H12) pressure on intramolecular C=C bond

formation (intra

r , ●) and intermolecular C=C bond formation (inter

r , ▲) in propanal (C3H6O) and

the rate ratio for intra- over intermolecular C=C bond formation (inter

r /intra

r , ○) during C3H6O

reactions on H-MFI catalysts at 473 K (Si/Al=11.5, 7.5 ks, 1.1×10−3

mol C3H6O· (mol H+

i·s)−1

,

1.1 kPa C3H6O in He).

Both the turnover rates for inter- and intramolecular C=C bond formation decreased as C6H12

pressure increased, because C6H12 species titrated a portion of the adsorbed propanal (and its

isomers), as expected from the higher proton affinity (~813 kJ·mol−1

for C6H12 2

vs. 786 kJ·mol−1

for C3H6O [40]) and higher dispersive interactions as they adsorb on MFI zeolites (increased by

15 kJ·mol−1

for each additional C atom in alcohol [38, 58] and 10-12 kJ·mol−1

for each

additional C atom in alkane [27, 59, 60]) for C6H12 than C3H6O. The rates for intramolecular

C=C bond formation (rintra), however, decreased to a much lower extent than those for

intermolecular C=C bond formation, because C6H12 promotes the kinetically relevant hydrogen

transfer step (Step 2.3, Scheme 2.2) but does not affect the nucleophilic addition step (Step 1.2,

Scheme 2.2).

0.00 0.05 0.10 0.150

5

10

15

20

25

30

rintra/rinter

rintra

rinter

r intr

a/r in

ter

Tu

rno

ve

r ra

te (

10

-5m

ol·

(mo

l H

+ r·s

)-1)

PC6H12 (kPa)

0.0

0.2

0.4

0.6

36

The rates of intermolecular C=C bond formation in C3H6O-C6H12 mixtures (3 6 6 12,C H O-C Hinter

r )

acquire an additional dependence on C6H12 pressure; thus Equation 2.2 becomes:

,

3 6

3 6 6 12

3 6 3 6 6 126 12

2taut C H Oaldol ads

C H O-C H

C H O C H O C Hads ads ads,C H

inter

taut,1

+ + +53 6 3 6 12

'

(C H O-H ) (C H OH-H ) (C H -H )

n

ss

k K K P

K P K K P K P

r

( ) ( ), ( ), and their

physisorbed forms

3 4a 4b

(2.10)

when H+ sites are predominantly occupied by protonated propanal and its isomers [(3), (4a), (4b),

and their physisorbed forms] and C6H12 as the most abundant surface intermediates. 6 12C HP and

6 12ads,C HK denote the partial pressure of C6H12 and the equilibrium constant for C6H12 adsorption

on H+ sites, respectively.

The rates of intramolecular C=C bond formation (3 6 6 12,C H O-C Hintra

r ) contain two distinct rate

constants for hydrogen transfer reactions with a portion of aromatics (predominantly

hydronaphthalenes) ( H transk ) and C6H12 (6 12H trans,C Hk ) as the H-donating agents, because C6H12

and hydronaphthalenes both participate as hydrogen donors but at different rates as a result of the

variation in C-H bond strength (e.g. 305-315 kJ∙mol−1

and 319-323 kJ∙mol−1

for the α-H of

hydronaphthalene and C6H12, respectively [57]):

2 6 12 6 12 3 6

3 6 6 12

3 6 6 126 12

,

H trans R'H H trans,C H C H C H Oads

C H O-C H

C H O C Hads ads,C H1

taut,1

intra

taut,

'

1 'n

s

n

ss

s

k P k P K K P

K K P K P

r

(2.11)

The term 2H trans R'Hk P corresponds to the rate of hydrogen donation from aromatics, as defined in

Equation 2.7, and the rest of the kinetic and thermodynamic parameters are defined in Sections

2.3.3 (Eqn. 2.2), 2.3.4 (Eqns. 2.6 and 2.7), and 2.3.6 (Eqn. 2.10). The total pressure of aromatics

varied from 4.5×10−4

to 7.0×10−4

kPa during propanal catalysis in C3H6O (without C6H12) (1.1-

4.5 kPa C3H6O, Fig. 2.5) and in C3H6O-C6H12 mixtures (0-0.15 kPa C6H12, 1.1 kPa C3H6O, Fig.

2.7). As a result, pressures of the portion of aromatics that act as hydrogen donors (2R'HP ) must

37

also remain at similar values despite the changes in C6H12 and C3H6O pressures during rate

measurements. Therefore, the term 2H trans R'Hk P in Equation 2.11 is approximated as a constant.

Combining Equations 2.11 and 2.10, the rate ratio for intramolecular to intermolecular C=C bond

formation (intra

r /inter

r ) is given by:

2 6 12 6 12

3 6

H trans R'H H trans,C H C Hintra

taut C H Ointer aldol

taut, 1

' ( )n

ss

K k P k Pr

r k K P

(2.12)

This expression predicts a linear increase in rintra/rinter with C6H12 pressure when the pressures of

propanal (3 6C H OP ) and hydrogen donors (

2R'HP , e.g. 6,8-dimethyl-1,2,3,4-tetrahydronaphthalene,

Scheme 2.1) were held relatively constant as the C6H12/C3H6O ratio was varied from 0 to 0.15, as

is shown in Figure 2.7. The selective promotion of C6H12 towards the hydrogen transfer step over

the nucleophilic addition step partially compensates for the reduction in propanal surface

coverage and leads, in turn, to an increase in the rate ratio for intramolecular over intermolecular

C=C bond formation, rintra/rinter, with increasing C6H12 pressure.

2.3.7. Regression of rate data with the derived rate expressions for inter- and intramolecular C=C bond formation

Rate equations derived for inter- and intramolecular C=C bond formation from Sections 2.3.3 to

2.3.6 (Eqns. 2.8, 2.9, 2.10, and 2.11) were used to regress against the rate data from C3H6O,

C3H6O-H2O, and C3H6O-C6H12 reactions over H-MFI at 473 K, presented in Figures 2.5, 2.6, 2.7,

and Figure S2.3 (in Appendix Sec. 2.6.4), to obtain the estimated kinetic and thermodynamic

parameters in Table 2.2.

38

Table 2.2. Rate parameters derived from non-linear regression fittings of rate data to rate

equations [Eqn. 2.8 (in Sec. 2.3.4), Eqn. 2.9 (in Sec. 2.3.5), Eqns. S2.2 and S2.3 (in Appendix

Sec. 2.6.4)]

Parameter Value

tautinter,eff aldol

taut,1

1 'n

ss

Kk k

K

a

1.97×10−4

s−1·kPa

−1 ± 0.06×10

−4 s

−1·kPa−1

6 10ads,C H O

-inter,eff aldol

taut,dehy ads1

(1 ' )n

ss

Kk k

K K K

a

9.5×10−4

s−1·kPa

−1 ± 1.8×10

−4 s

−1·kPa−1

2

taut,1

H trans R'Hintra,eff

taut,1

'

1 '

n

ss

n

ss

K

k k P

K

b

4.18×10−5

s−1

± 0.10×10−5

s−1

6 12ads,C H

taut, ads1

1 'n

ss

K

K K

c

15.9 ± 1.2

6 12

2

H trans, C H

H trans R'H

k

k P c

4.3 kPa−1

± 0.6 kPa−1

a Estimated values for kinetic parameters in Equation 2.9 (Sec. 2.3.5); see Appendix Section 2.6.4 for the determination of

inter,effk and -inter,effk .

b Estimated values for kinetic parameters in Equation 2.8 (Sec. 2.3.4). c Estimated values for kinetic parameters in Equations S2.2 and S2.3 (in Appendix Sec. 2.6.4).

The effective rate constant for intermolecular C=C bond formation (kinter, eff) is:

tauttautinter,eff aldol aldol

taut,1

1 'n

ss

Kk k k K

K

for taut,

1

1 'n

ss

K

(2.13)

39

and its value was found to be 1.97×10−4

±0.06×10−4

s−1

·kPa−1

at 473 K. The two terms in the

denominator of Equation 2.13, 1 and taut,1

'n

ss

K , correspond to the coverages of the protonated

propanal [(3), Scheme 2.2] and adsorbed propanal isomers [propenol (4a), allyl alcohol (4b), and

their physisorbed forms], respectively. Protonated propenal (4a) and allyl alcohol (4b) were

estimated to be 94 kJ·mol─1

and 71 kJ·mol─1

less stable than protonated propanal (3) in H-MFI

zeolites, based on their relative proton affinities with ammonia and on the heat of ammonia

adsorption but not accounting for the difference in dispersive interaction energies between these

species and the MFI pore walls [33]. Similar trends and values are expected for the physisorbed

species because of the small energy differences between the protonated and physisorbed species

[36-39]. The relative magnitudes of these heats of adsorption suggest that the second term

taut,1

'n

ss

K is much smaller than 1, thus inter,eff

k is approximated to be equal to tautaldolk K .

The rate constant for the reverse reaction (reverse of Step 1.2, Scheme 2.2), inter,effk , for

intermolecular C=C bond formation was found to be 9.5×10−4

±1.8×10−6

s−1

·kPa−1

, a magnitude

similar to the forward rate constant (1.97×10−4

s−1

·kPa−1

±0.06×10−4

s−1

·kPa−1

) at 473 K. This

reaction becomes reversible as the conversion increases and in the limit of high2H OP /

3 6C H OP (e.g.

reverse,interr / net,inter

r =0.45 when 2H OP /

3 6C H OP =9, 473 K) or high 6 10C H OP /

3 6C H OP (not shown here)

ratios. The reverse rates are unimportant and do not affect the net rates at the low C6H10O and

H2O pressures used in the rate measurements for obtaining the rate data reported in Figure 2.6

(6 10C H OP /

3 6C H OP = 0.005-0.01).

Turnover rates for the intramolecular C=C bond formation equal the effective rate constant

intra,effk (Eqn. 2.8) and depend on the aggregate values of the elementary rate constants for

hydrogen transfer ( H transk ) and the pressures of hydrogen donors (2R'HP ), as defined in Equation

2.7, and the sum of equilibrium constants for propanal adsorption at various conformations

( taut,1

'n

ss

K , as defined in Sec. 2.3.3, Eqn. 2.2). These effects of pressures and rate constants

40

were lumped and treated here as a pseudo rate constant, because the pressures of the hydrogen

donors were relatively constant during rate measurements (Sec. 2.3.6):

2H trans R'H

taut,1

intra,eff

taut,1

'

1 '

n

ss

n

ss

K

k k P

K

(2.14)

These assumptions (Eqns. 2.13 and 2.14), upon substitution into Equation 2.12, give the rate

ratio for intermolecular over intramolecular C=C bond formation, which reflects the selectivity

ratio for C6H10O (2-methyl-2-pentenal) to C3H6 formation in the primary reaction paths

(6 10

3 6

C H OC H

S ):

3 6

6 10

3 6

6 12 6 12

C H OC H

2

C H Oaldol

H trans R'H H trans,C H C H

tautinter

intrataut,

1

'n

ss

k K PrS

rK k P k P

(2.15)

The selectivities towards C6H10O, according to Equation 2.15, are expected to increase with

increasing propanal pressure (Fig. 2.5), because higher propanal pressure favors the

intermolecular over intramolecular C=C bond formation.

Non-linear regression of the rate data measured in C3H6O-C6H12 mixtures (Fig. 2.7) against

Equations S2.2 and S2.3 (rearranged in simplified forms from Eqns. 2.10 and 2.11, respectively,

see Appendix Sec. 2.6.4) gives the rate parameters for hydrogen transfer (Table 2.2). The term

26 12

H trans R'HH trans,

1

C Hk k P

(from Eqn. S2.3 in Sec. 2.6.4) represents the relative reactivities of

C6H12 to hydronaphthalene species as hydrogen donors; its value was found to be 4.3±0.8 kPa−1

(473 K). The rate data reported in Figure 2.7 were measured at aromatic pressures of less than

10−3

kPa and within the aromatic lump only a small portion of them acting as hydrogen donors.

Substituting 10−3

kPa as the maximum value of hydrogen donor pressure (as 2R'HP ) gives the

6 12

1

H transH trans,C Hk k

ratio of 4.3×10−3

. The 6 12

1

H transH trans,C Hk k

ratio much lower than unity

is consistent with much lower reactivities for hydrogen transfer in C6H12 than in

hydronaphthalene species, as predicted from the differences in their C-H bond strength

41

[estimated to be 319-323 kJ∙mol−1

for the tertiary allylic C-H bond in C6H12 [57] vs. 305-315

kJ∙mol−1

for the H leaving group in hydronaphthalene [57]]. The ratio

6 12,C H taut,ads ads1

1

1 'n

ss

K K K

was estimated to be 16.5±1.7, a value larger than unity and

indicates a higher heat of adsorption for C6H12 than those for propanal and its isomers [(3), (4a),

(4b) and their physisorbed forms, Scheme 2.2] on H-MFI zeolite, consistent with the trend in

their proton affinity and dispersive interaction energy (Sec. 2.3.6). These studies highlight the

key requirements that determine the catalytic fate of propanal to either lengthen or preserve its

carbon chain length during its catalytic sojourns on H+ sites. These mechanistic insights would

allow us to quantitatively predict and control rates and product distributions.

2.4. Conclusion

Kinetic interrogations and acid site titrations lead to a proposed sequence of elementary steps for

propanal deoxygenation on H+ sites immobilized within MFI framework. The reaction occurs via

competitive pathways of inter- and intramolecular C=C bond formation that evolve 2-methyl-2-

pentenal and propene, respectively. The intermolecular C=C form formation occurs via coupling

of propanal and intramolecular C=C bond formation via direct oxygen removal as H2O. These

reactions proceed in parallel with distinct rate dependencies on Brønsted acid sites

predominantly occupied by protonated propanal and its isomers, which present as a kinetically

indistinguishable lump during steady-state catalysis. Rates for intermolecular propanal coupling

increase linearly with propanal pressure and decrease with water pressure and are limited by the

nucleophilic attack of propenol to protonated propanal in a bimolecular condensation reaction.

The rates of intramolecular C=C bond formation, however, remain insensitive to propanal

pressure, because rates are limited solely by the transfer of hydrogen atom from hydrogen

donating agents to propanal derived surface intermediates. Water as a byproduct mitigates site

occupation by larger, inactive carbonaceous species and increases the reverse rates for

intermolecular C=C bond formation, thus decreasing the net rates of this step. Water, however,

does not affect the net rates for intramolecular C=C bond formation because this step is

irreversible. The kinetic relevance of hydrogen transfer in intramolecular C=C formation step is

confirmed from the increase in the rate ratio for intra- over intermolecular C=C bond formation

42

as the pressure of hydrogen donating agent increases. During steady-state catalysis, hydrogen

donating events, enabled by sequential reactions of intramolecular ring closure and

dehydrogenation steps that remove hydrogen and increase the degree of unsaturation in

secondary products, and the hydrogen accepting events are kinetically coupled to evolve propene

as a primary product. This knowledge on the nature of surface intermediates and kinetic

requirements in catalytic rates and selectivities provide the framework on tuning the relative

rates for the initial inter- and intramolecular C=C bond formation in propanal and the product

distributions during propanal deoxygenation on Brønsted acid sites contained within the MFI

structures.

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2.6. Appendix

2.6.1. Mass balance during propanal reaction on H-MFI

At 2.1 ks, the carbon input into the reactor (1.2×10−6

mol C·s−1

) is much higher than output

(0.87×10−6

mol C·s−1

), indicating a lot of coke or other heavy compounds (e.g. naphthalenes

whose boiling points are over 473 K) are produced and condensed on the catalysts. After 3.9 ks,

the catalysts have been deactivated due to the coverage of the heavy products (e.g. naphthalenes

whose boiling points are over 473 K), as a result, the propanal conversion decreases to below

5 %, which is a differential condition. Unfortunately, limited by the accuracy of the feeding

system which typically has a fluctuation of ±5%, it is hard to report the accurate mass balance of

the reaction under differential condition. However, when I raised the temperature to 673 K, the

conversion was over 90 % and the carbon input and output in the stream were found to be

balanced as shown in Figure S2.1.

47

Figure S2.1. Evolution of carbon in the products (○, excluding C3H6O) and unreacted C3H6O (□)

in the reactor effluent stream, total carbon in the reactor effluent stream (▲), and total carbon in

the feed mixture (dash line) as a function of time-on-stream during propanal (C3H6O) reactions

on H-MFI at 673 K (Si/Al=11.5, 5.3×10−2

mol C3H6O·(mol H+

i·s)−1


2.6.2. Time on stream evolution of propanal conversion

For reaction times below 2.1 ks, propanal conversions exceeded 30 % (see Fig. S2.2 on the time-

dependent conversion and selectivity values) and lower amount of total carbon species in the

product stream compared with that of the feed stream suggest the accumulation of heavier

products inside the zeolitic pores. For reaction times above 7.5 ks, propanal conversions were

lower (< 5%), the rates of carbon accumulation within the zeolitic pores were negligible (Fig.

S2.1), and conversion and selectivity values remained unchanged during rate measurements. The

rate values for inter- and intramolecular C=C bond formation (inter

r and intra

r ) at 7.5 ks, which

reflect the steady-state catalytic rates, are plotted as a function of C3H6O pressure in Figure 2.5

and used here for kinetic analysis.

0 5 10 150

5

10

15

20

25

Total input carbonTotal output carbon

Products

Ca

rbo

n in

th

e r

ea

cto

r e

fflu

en

t str

ea

m (m

ol

s1g

ca

t.

1)

Time on stream (ks)

Unreacted C3H6O

48

Figure S2.2. Propanal (C3H6O) conversion (■) and selectivities towards intermolecular C=C

bond formation (Sinter, ▲) and intramolecular C=C bond formation (Sintra, ●) as a function of

time-on-stream during propanal reactions on H-MFI at 473 K (Si/Al=11.5, 1.1×10−3

mol

C3H6O·(mol H+

i·s)−1


2.6.3. Determination of kinter,eff and k−inter,eff in Equation 2.9

Propanal reactions with a C3H6O-H2O mixture were carried out at constant propanal pressure

(3 6C H OP = 1.1 kPa), so term

3 6C H Ointer,effk P in Equation 2.9 remains a constant value. This

condition simplifies the rate of intermolecular C=C bond formation (inter

r , Eqn. 2.9) to:

6 10 2

3 6

3 6

C H O H O

C H O

C H Ointer inter,eff inter,eff

P Pr k P k

P (S2.1)

Rates for intermolecular C=C bond formation (inter

r ) were plotted against 6 10 2 3 6C H O H O C H O

1P P P

(Fig. S2.3) and linear regression was carried out to obtain the slope inter,effk

= ─9.5×10−4

s−1·kPa

−1 ± 1.8×10

−4 s

−1·kPa−1

, and the intercept 3 6C H Ointer,eff

k P = 2.17×10−4

s−1·kPa

−1 ± 0.07×10

−4

s−1·kPa

−1, for Equation S2.1, these values of inter,eff

k and inter,effk are calculated and summarized

in Table 2.2.

0 2 4 6 80

10

20

30

40

Sintra

Se

lectivitie

s (

%)

Pro

pa

na

l co

nve

rsio

n (

%)

Time on stream (ks)

Sinter

0

20

40

60

80

100

49

Figure S2.3. Effects of H2O on rinter during propanal (C3H6O) reactions on the H-MFI

(Si/Al=11.5) at 473 K (7.5 ks, 1.1×10−3

mol C3H6O·(mol H+

i·s)−1


2.6.4. Determination of kinetic parameters in Equations 2.10 and 2.11 by non-linear regression fitting of these equations with rate data for C3H6O-C6H12 reactions on H-MFI zeolites in Figure 2.7

Rate expressions for inter- and intramolecular C=C bond formation in C3H6O-C6H12 feed

mixtures, Equation 2.10 and Equation 2.11, discussed in Section 2.3.6, were re-arranged to

Equations S2.2 and S2.3, respectively:

6 12

3 6

3 6

3 6 6 126 12

6 126 12

3 63 6

tautaldolC H O

C H O

, C H O-C H,C H ,C H

C HC H

C H OC H O

taut,inter, eff1

ads ads

taut,taut, adsads11

inter

1 '

11

1 '1 '

n

ss

nn

ssss

k KP

K k P

K KPP

K K PK K P

r

(S2.2)

where aldol tautinter,eff

taut,1

1 'n

ss

k Kk

K

as shown in Equations 2.3 and 2.9;

0.00 0.02 0.04 0.06 0.080

1

2

3

4

r in

ter

(10

-4m

ol·

(mo

l H

+ r·s

)-1)

PC6H10OPH2OP

C3H6O

1(kPa)

50

2 6 12 6 12

6 12

2

3 6 6 126 126 12

3 6

H trans R'H H trans,C H C H

H trans,C H

H trans R'H

,C H O-C HC Hads,C H

C H O

taut, taut,1 1

intra,efftaut, taut,1 1

taut,ads1

intra

' '

11 ' 1 '

1

1 '

n n

s ss s

n n

s ss s

n

ss

K k P K k P

kkK K k P

K P

K P K

r

6 12

6 12

6 12

3 6

C H

,C H

C H

C H O

ads

taut,ads1

1

1 'n

ss

P

KP

K P K

(S2.3)

where 2H trans R'H

taut,1

intra,eff

taut,1

'

1 '

n

ss

n

ss

K

k k P

K

as in Equations 2.8 and 2.14.

Propanal pressures C H O3 6P were maintained at constant values of 1.1 kPa and treated as a

constant in rate data fittings. 2H trans R'Hk P was also treated as a constant, because the pressures of

aromatics that act as hydrogen donors (2R'HP ) remain relatively stable, as discussed in Section

2.3.6. Equations S2.2 and S2.3 were simplified to:

3 6 6 12

6 12

C H O

C H

C H O

3 6

3 6

, C H O-C H

inter,eff

inter

1

k P

aP

P

r

(S2.4)

6 12

3 6 6 12

6 12

C H

C H

C H O3 6

, C H O-C H

intra,eff

intra

1

1

k b P

aP

P

r

(S2.5)

where 6 12,ads C H

taut, ads1

1 'n

ss

Ka

K K

; 6 12

2

H trans,C H

H trans R'H

kb

k P .

By substituting the values of inter,effk (Table 2.2) and C H O3 6

P (1.1 kPa) into Equation S2.4, and

carrying out non-linear regression fitting with the rate data of intermolecular C=C bond

formation ( interr ) in Figure 2.7, the value of a was obtained as shown Table S2.1 and Table 2.2.

Afterward, the values of intra,effk (Table 2.2), C H O3 6

P (1.1 kPa) and a (Table S2.1) were

substituted into Equation S2.5, followed by the non-linear regression with the rate data of

51

intramolecular C=C bond formation ( intrar ) data in Figure 2.7. The value of b was obtained as in

Table S2.1, as well as in Table 2.2.

Table S2.1. Parameter values derived from non-linear regression of rate data in Figure 2.7 (H-

MFI, Si/Al=11.5, 7.5 ks, 1.1×10−3

mol C3H6O·(mol H+

i·s)−1

, 1.1 kPa C3H6O in He, 473K) with

Equations S2.4 and S2.5

Parameter Value

6 12ads,C H

taut, ads1

1 'n

ss

Ka

K K

15.9 ± 1.2

6 12

2

H trans,C H

H trans R'H

kb

k P

4.3kPa−1

± 0.6 kPa−1

52

Chapter 3 Alkanal Transfer Hydrogenation Catalyzed by Solid Brønsted

Acid Sites

Abstract

Catalytic pathway and requirements for transfer hydrogenation of n-alkanals (CnH2nO, n=3-6) on

Brønsted acid sites (H+) immobilized in microporous MFI and FAU crystalline structures or

dispersed on H4SiW12O40 polyoxometalate clusters are established by isolating its rates from

those of the various concomitant catalytic cycles. Transfer hydrogenation of alkanals involves a

kinetically relevant, intermolecular hydride transfer step from substituted tetralins or

cyclohexadienes produced from the parallel alkanal coupling and ring closure reactions as the

hydride donor (R’H2) to protonated alkanals (RCH2CHOH+) as the hydride acceptor, via a

bimolecular transition state with a shared hydride ion, (RCH2CHOH+∙∙∙H

−∙∙∙R’H

+)‡. The rate

constants for the intermolecular hydride transfer step correlate directly to the hydride ion affinity

difference between the carbenium ions of the H-donors (R’H+) and the protonated alkanals

(RCH2CHOH+). As a result, smaller alkanals with higher hydride ion affinities are more

effective in abstracting hydride ions and in transfer hydrogenation (C4>C5>C6). Propanal is an

exception, as it is less effective in transfer hydrogenation than butanal. The deviation of propanal

from the reactivity trend is apparently caused by its smaller transition state for hydride transfer,

which is solvated to a lesser extent in FAU cages. The transfer hydrogenation occurs much more

effectively on partially confined H+ sites in FAU structures than in smaller pore MFI or

unconfined H4SiW12O40 polyoxometalate clusters, an indication that FAU solvates and stabilizes

the bulky transition state of hydride transfer via van der Waals interactions. These effects of local

site structures and the thermochemical properties of reactant determine the reactivity of alkanal

transfer hydrogenation and thus selectivity ratio of alkenes, dienes, aromatics, and larger

oxygenates during deoxygenation catalysis.

53

3.1. Introduction

Brønsted acid sites (H+) immobilized on solid matrixes catalyze the deoxygenation of light

alkanals (RCH2CHO, R=CH3, C2H5, or C3H7) in steps that involve intermolecular or

intramolecular C=C bond formation, isomerization, dehydration, ring closure, and

transalkylation reactions and form alkenes, dienes, alkenals, and aromatics at moderate

temperatures (473-673 K) and ambient pressure, as established on H-MFI [1-4], H-Y [5], and

H4SiW12O40 polyoxometalate clusters [6]. Contained within these concurrent catalytic steps is

the direct alkanal deoxygenation, which converts an alkanal reactant to the corresponding alkene

(RCH2CHO+2HRCHCH2+H2O) [1]. This reaction, in the absence of external hydrogen

sources, must involve intermolecular shuffling of hydrogen from reaction products to alkanal

reactants, as required by the reaction stoichiometry. Despite the obvious involvement of reaction

products as the hydrogen donors in these ubiquitous transfer hydrogenation events, their

mechanism and site requirements have not been clearly established.

Brønsted acid catalyzed hydride transfer has been studied extensively with density functional

theory (DFT) calculations for the transfer from alkanes to alkenes on H3Si-OH-AlH2-O-SiH3

clusters [7-9], from alkanes (e.g., propane and t-butane) to alkoxides (e.g., propyl and t-butyl

alkoxides) in mordenite zeolite [10], and from alkanes (e.g., methane and ethane) to their

corresponding carbenium ions (e.g., methyl and ethyl carbenium ions) in the gas phase [11]. It

has also been probed experimentally between isobutane and cyclohexene on beta and ZSM-5

zeolites and on sulfated zirconia [12], during alkane cracking in zeolites (e.g., SAPO-41, ZSM-5,

and Y) [13], and during dimethyl ether homologation on H4SiW12O40 cluster, FAU zeolite, and

mesoporous SiO2-Al2O3 catalysts [14]. Hydride transfer on Brønsted acid occurs when a hydride

ion donor (H-donor) donates a hydride ion to the hydride ion acceptor (H-acceptor) via the

formation of a carbonium ion transition state sharing the hydride ion. An example of the H-donor

is an alkane and of H-acceptor is either an adsorbed carbenium ion at the H+ site [7, 8, 11, 15] or

an alkoxide [9, 10, 15] at the ground state [16-18]. The transition state decomposes when the H-

acceptor desorbs while the H-donor becomes a carbenium ion and then donates its proton back to

the catalyst surfaces, thus regenerating the H+ site and completing the catalytic cycle [7, 8].

Local confinement of the H+ site appears to influence the hydride transfer reactivity [13, 14]:

hydride transfer from n-C5H12 to C2H5+, C3H7

+, and C4H9

+ carbenium ions is most effective when

occurring inside zeolites with cavity volumes of ~0.2 nm3 [e.g., CIT-1 (0.211 nm

3) and MCM-68

54

(0.182 nm3)] than those with either larger [e.g., MCM-22 (0.467 nm

3) and Y (0.731 nm

3)] or

smaller [e.g., SAPO-41 (0.081 nm3) and ZSM-5 (0.131 nm

3)] cavity volumes [13]. The reaction

apparently requires its bimolecular transition state to be at comparable dimensions with those of

the zeolite cavities. In fact, the sums of the volumes for (C2H5+-n-C5H10), (C3H7

+-n-C5H10), and

(C4H9+-n-C5H10) fragments are estimated to be 0.182, 0.203, and 0.224 nm

3, respectively, in

similar magnitudes with the cavities in CIT-1 and MCM-68 zeolites [13].

Large-pore FAU and BEA zeolites exhibit higher selectivities towards triptane than mesoporous

SiO2-Al2O3 and medium-pore MFI zeolites during solid acid-catalyzed dimethyl ether

homologation [14], because confinement within the larger pores preferentially solvates the larger

transition states for hydride transfer and methylation and terminates the chain growth at C7

products (triptane). Despite these extensive studies on the hydride transfer between alkanes and

alkenes, few studies have addressed the hydride transfer to protonated carbonyl species. A recent

study on hydrogen transfer and sequential dehydration of naphthols on H-Y zeolites has reported

an increase in hydrogen transfer rates in the presence of hydrocarbons (e.g., tetralin and 1,5-

dimethyltetralin), as these hydrocarbons may act as the hydrogen donors [19]. It is hypothesized

that a hydride ion is being transferred from the H-donor to the keto tautomer of naphthol and the

hydride ion dissociation energy of the H-donor influences the rate [19]. Little mechanistic details

are available for the transfer hydrogenation of n-alkanals, despite the clear kinetic evidence of

their predominant occurrences during their deoxygenation on solid Brønsted acid catalysts.

Here, I report catalytic insights and kinetic requirements for the transfer hydrogenation events,

which shuffle hydrogen from H-donors, identified to be aromatic species (e.g., alkyl tetralins) or

precursors to aromatics (e.g. alkyl cyclohexadienes), to protonated alkanals at Brønsted acid sites

(H+) in MFI and FAU zeolites or on polyoxometalate (H4SiW12O40) clusters. I show that transfer

hydrogenation occurs in a direct, concerted step between substituted tetralins or alkyl

cyclohexadienes and protonated alkanals. The hydride donors are products of intermolecular

C=C bond formation, ring closure, and dehydrogenation reactions. The hydride transfer

reactivity exhibits a clear correlation with the hydride ion affinity differences between the

carbenium ions of the H-donor and the H-acceptor and is a strong function of the extent of local

structural confinements around the H+ sites.

55

3.2. Experimental


H-MFI and H-FAU zeolite samples were prepared by treating their NH4+ form (Zeolyst,

CBV2314, 425 m2 g

−1, Si/Al atomic ratio=11.5, Na2O=0.05 wt.%) and H

+ form (Zeolyst,

CBV720, 780 m2 g

−1, Si/Al atomic ratio=15, Na2O=0.03 wt.%), respectively, in flowing dry air

(Linde, zero grade, 0.6 cm3 gcat.

−1 s

−1), by heating to 873 K at 0.0167 K s

−1 and then holding

isothermally at 873 K for 4 h. H4SiW12O40/SiO2 catalysts (0.075 mmol H4SiW12O40 gSiO2−1

) were

prepared by dispersing H4SiW12O40 (Sigma Aldrich, reagent grade, CAS #12027-43-9) on

chromatographic SiO2 (GRACE, 330 m2 g

−1, 0-75 μm, 1.2 cm

3 g

−1 pore volume, treated in air at

673 K for 5 h) via incipient wetness impregnation with a solution of H4SiW12O40 and ethanol

(Sigma-Aldrich, >99.5%, anhydrous). The impregnated H4SiW12O40/SiO2 samples were held in

closed vials for 24 h and then treated in flowing dry air (Linde, zero grade, 0.1 cm3 g

−1 s

−1) at

323 K (0.0167 K s−1

heating rate) for 24 h. The H+ site densities on these catalysts (mol H

+ gcat.

−1)

were measured by pyridine titration at 473 K, as described in our previous work [1].

3.2.2 Rate and selectivity assessments

Alkanal conversion rates and site-time-yields of alkenes, dienes, oxygenates, and aromatics were

measured in a fixed bed microcatalytic quartz reactor (9.5 mm inner diameter), which was loaded

with 100 mg H-MFI or H-FAU zeolites or 50 mg H4SiW12O40/SiO2 powders supported on a

coarse quartz frit. Catalysts were treated in-situ under flowing He (Linde, Grade 5.0, 8.3-16.7

cm3 gcat.

−1 s

−1) at 0.0167 K s

−1 to the reaction temperature (573 K) prior to rate measurements.

Alkanal or butanol reactants [butanal (Sigma Aldrich, puriss grade, ≥99%, CAS# 123-72-8),

propanal (Sigma Aldrich, 97%, CAS# 123-38-6), pentanal (Sigma Aldrich, 97%, CAS# 110-62-

3), hexanal (Sigma Aldrich, 98%, CAS# 66-25-1), or butanol (Sigma Aldrich, 99 %, CAS# 71-

36-3)] were introduced via a gas tight syringe (either 5 cm3

Hamilton Model 1005 or 1 cm3 SGE

Model 008025), which was mounted on a syringe infusion pump (KD Scientific, LEGATO 100),

into a vaporization zone heated to the boiling points of the respective reactants at atmospheric

pressure, within which liquid alkanals were evaporated and mixed with a He (Linde, Grade 5.0,

8.3-16.7 cm3 gcat.

−1 s

−1) or H2 (Linde, Grade 5.0, 8.3 cm

3 gcat.

−1 s

−1) purge stream. The mixture

was fed to the reactor via heated transfer lines held isothermally at 473 K. Tetralin (Sigma

Aldrich, 99%, CAS# 119-64-2), tetralin-adamantane mixture with a molar ratio of 20:1

56

(adamantane, Sigma Aldrich, 99%, CAS# 281-23-2), or cyclohexadiene (Sigma Aldrich, 97 %,

CAS# 592-57-4) was introduced into a second vaporization zone, which was located downstream

of the zone for alkanal or alkanol vaporization described above, through a gas tight syringe (0.25

cm3 SGE Model 006230) mounted on a syringe infusion pump (KD Scientific, LEGATO 100).

This vaporization zone was maintained at 458 K for tetralin or tetralin-adamantane mixture

infusion and 353 K for cyclohexadiene infusion. Chemical species in the reactor effluent stream

were quantified with an on-line gas chromatograph (Agilent, Model 7890A) and mass

spectrometer (Agilent, Model 5975C) by chromatographic separation with HP-5 (Agilent,

19091J-413, 30 m, 0.32 mm ID) or HP-5MS (Agilent, 190091S-433, 30 m, 0.25 mm ID)

capillary columns. The HP-5 column was connected to thermal conductivity (TCD) and flame

ionization (FID) detectors installed in series and the HP-5MS column to the mass spectrometer

(MS). For each data point, the carbon balance, defined by the difference between the molar flow

rates of all carbon species contained in the feed and the reactor effluent stream, was less than

10 %.


3.3.1. Alkanal deoxygenation pathways and the kinetic couplings of intramolecular C=C bond formation in alkanals and dehydrogenation of aromatic products at Brønsted acid sites

Catalytic pathways for alkanal (propanal [1, 2] and butanal [6]) deoxygenation on solid Brønsted

acid sites (H-MFI [1, 2] and H4SiW12O40 [6]) shown in Scheme 3.1 have been previously

established based on selectivity changes with residence time and confirmed from reactions with

the intermediates [2, 6]. Butanal (C4H8O) deoxygenation occurs on Brønsted acid sites (H+) via a

bimolecular, aldol condensation-dehydration step (Step 1, Scheme 3.1), which creates an

intermolecular C=C bond and forms 2-ethyl-2-hexenal (C8H14O), before its successive reactions

with another butanal (Step 1.2) to evolve 2,4-diethyl-2,4-octadienal (C12H20O). These larger

alkenals (including 2-ethyl-2-hexenal and 2,4-diethyl-2,4-octadienal) undergo sequential

cyclization-dehydration (Steps 1.1.1 and 1.3.1) or cyclization-dehydration-dehydrogenation

(Step 1.3.2) reaction that forms cycloalkadienes or aromatic species. Dehydrogenation of the

57

cycloalkadienes and substituted tetralin species via Steps 1.1.2 and 1.3.3, respectively, increases

their extents of unsaturation and leads to substituted benzenes (e.g., xylene) and naphthalenes

(e.g., 1,3-dimethylnaphthalene), respectively, which upon transalkylation reactions evolve

diverse aromatics (C7-C19, not shown in Scheme 3.1) [20, 21]. Butanal may also undergo a

primary, intramolecular C=C bond formation, during which it accepts two H atoms followed by

dehydration to evolve butene (Steps 2.1-2.2). An alternative, competitive isomerization-

dehydration (Steps 3.1-3.2) of butanal leads to butadiene [22]. A small number of basic sites in

H-MFI and H-FAU zeolites (e.g., 0.05 and 0.03 wt.% Na2O in H-MFI and H-FAU, respectively)

catalyze Tishchenko esterification reaction (Step 4.1), which transforms two butanals into butyl-

butyrate and sequential ketonization (Step 4.2) and hydrogenation-dehydration (Step 4.3) evolve

3-heptene [5]. The rates of Tishchenko esterification reaction (Step 4.1) increase proportionally

with the number of basic sites (Fig. S3.1, Appendix), which include the bi-coordinated oxygen in

the extra-framework alumina [23] and, for Na-exchanged H-MFI zeolites, at the conjugated

oxygen of Na+ ions [24]. Thus, this reaction occurs strictly at the basic sites. These basic sites,

however, are essentially inactive for intermolecular C=C bond formation (Step 1) at 573 K, as

confirmed from the proportional decrease in its rates with the H+ site density. As the number of

H+ sites goes to zero, the rates for intermolecular C=C bond formation approach zero as well,

despite the increase in the basic site (Na+) density on a series of Na-exchanged H-MFI zeolites

(Fig. S3.1a, Appendix).

These reactions occur in sequence or parallel on H-MFI, H-FAU, and H4SiW12O40/SiO2 catalysts.

Their rates and carbon selectivities are denoted as rj,m and Sj,m, respectively, where subscript j

represents the identity of reaction pathway (j=Inter, Intra, Dehy, or Tish, which denote inter- or


ketonization, respectively) and m represents the reactant (e.g., m=C4H8O). The rates of butanal

conversion (4 8,C H Ojr ) and selectivities (

4 8,C H OjS ) to various pathways j on H-MFI, H-FAU, and

H4SiW12O40 were measured at 573 K and the amounts of H+ sites remaining after the reaction at

different time-on-streams were determined by chemical titration with pyridine (as shown in Fig.

S3.2, Appendix). During the initial 125 min, butanal conversion rates (denoted as 4 8overall,C H Or ) on

H-MFI and H-FAU decreased by >72% and >47%, respectively, and the carbon selectivities

(4 8,C H OjS ) commensurately changed (Figs. S3.3a and S3.3b, Appendix), because of (1) the

58

gradual occupation of the H+ sites by butanal and (2) the loss of H

+ site (Figs. S3.2a and S3.2b,

Appendix) caused by the formation of heavier products (e.g. larger aromatics and coke) inside

the zeolitic pores. The rates of change for both the 4 8,C H Ooverallr and

4 8,C H OjS became significantly

smaller above 125 min. Above 125 min, the changes in rate per unit time, defined as

4 8,C H Ooverallr (∆time-on-stream) −1

, were one order of magnitude smaller than the initial values

and the changes in selectivity 4 8,C H OjS were less than ±6% over the course of 240 min for H-MFI

and H-FAU (Figs. S3.3a and S3.3b, Appendix). Similarly, the number of active H+ sites, butanal

conversion rates, and carbon selectivities on H4SiW12O40 became stable above 125 min (Figs.

S3.2c and S3.3c, Appendix). Based on these time-dependent results, it is concluded that butanal

reactions on all three catalysts reached steady-state after 125 min. The overall butanal conversion

rates, together with the rates and selectivities for each primary pathway (Pathways 1, 2, 3, and 4

in Scheme 3.1) at 125 min are summarized in Table 3.1. In the following, the rate of a primary

pathway was determined from the concentration of the primary product as well as those of the

secondary products resulting from the sequential reactions.

59

Scheme 3.1. Reaction network for butanal deoxygenation on solid Brønsted acid catalysts (“D”

and “A” denote H-donor and H-acceptor, respectively; most of the intermediates and products

shown in the scheme were detected in the experiment except crotyl alcohol and butanol because

of their rapid dehydration).

60

Table 3.1. Rates and selectivities for butanal deoxygenation and butanol dehydration on H-MFI,

H-FAU, or H4SiW12O40 at 573 K

Reactant Reaction Rate

(Carbon selectivity)

H-FAUf H-MFI

g H4SiW12O40

h

The rates are given in 10−2

mmol (mol H+ s)

−1

Butanal 4 8,C H Ooverallr b

(Overall conversion)

106

(14.3 %)

35.0

(10.5 %)

1530

(18.2 %)

Butanal Pathway 1a:

2C4H8O C8H14O+H2O

4 8,C H OInterr b

(4 8,C H OInterS c

)

15.2

(0.29)

6.6

(0.38)

652

(0.85)

Butanal Pathway 2a:

C4H8O +2H C4H8+H2O

4 8,C H OIntrar b

(4 8,C H OIntraS c

)

36.0

(0.37)

9.6

(0.27)

10.0

(0.01)

Butanal Pathway 3a:

C4H8O C4H6+H2O

4 8,C H ODehyr b

(4 8,C H ODehyS c

)

1.0

(0.01)

4.5

(0.13)

29.2

(0.02)

Butanal Pathway 4a:

2C4H8O C7H14+CO+H2O

4 8,C H OTishr b, d

in 10−5

mmol(gcat. s) −1

(4 8,C H OTishS c

)

7.4

(0.27)

0.9

(0.05)

1.2

(0.01)

1-Butanol C4H9OH C4H8+ H2O 4 9,C H OHDehyr e

- >370 >4400

Rate ratio

4 9 4 8

1

,C H OH ,C H ODehy Intra( )r r

-

>38

>440

aPathways 1, 2, 3, and 4 denote the reaction pathways (in Scheme 3.1) of intermolecular C=C bond formation (Step 1),

intramolecular C=C bond formation (Steps 2.1-2.2), isomerization-dehydration (Steps 3.1-3.2), and Tishchenko esterification-

ketonization (Steps 4.1-4.3), respectively;

b

4 8,C H Ooverallr denotes the overall C4H8O conversion rate;

4 8,C H Ojr represents the rate of reaction for pathway j during C4H8O

deoxygenation (subscript j=Inter, Intra, Dehy, or Tish, which denote inter- or intramolecular C=C bond formation, isomerization-

dehydration, or Tishchenko esterification-ketonization, respectively); c

4 8,C H OjS represents the selectivity to pathway j during C4H8O deoxygenation, defined as the rate of C4H8O consumption in

reaction j divided by the overall C4H8O conversion rate (subscript j=Inter, Intra, Dehy, Tish, which denote inter- or

intramolecular C=C bond formation, isomerization-dehydration, or Tishchenko esterification-ketonization, respectively);

dBecause Pathway 4 occurs on the basic sites, the unit for 4 8,C H OTish

r is given in terms of 10−5 mmol (gcat. s) −1, cat.=H-MFI, H-FAU,

or H4SiW12O40;

eThe rate of 1-butanol (C4H9OH) dehydration, 4 9,C H OHDehy

r , was measured with 1.1 kPa 1-butanol, space velocity 0.0033 and 0.045

mol 1-butanol (mol H+ s) −1 on H-MFI and H4SiW12O40, respectively; fSi/Al=11.5, space velocity 0.0033 mol butanal (mol H+ s) −1, 1.1 kPa butanal, time-on-stream=125 min; gSi/Al=15, space velocity 0.0074 mol butanal (mol H+ s) −1, 1.1 kPa butanal, time-on-stream=125 min;

61

h0.075 mmol H4SiW12O40 gSiO2−1, space velocity 0.045 mol butanal (mol H+ s) −1, 1.1 kPa butanal, time-on-stream=125 min.

These catalysts show different selectivity values towards the different paths (Table 3.1). H+ sites

on H4SiW12O40 clusters preferentially catalyze the intermolecular C=C bond formation (Pathway

1) with a carbon selectivity of 85 %, whereas H-FAU and H-MFI zeolites favor the

intramolecular C=C bond formation (Pathway 2) and isomerization-dehydration (Pathway 3)

reactions, both of which involve the catalytic sojourn of a single alkanal. This is caused in a large

part by the difference in the extent of H+ site confinement (to be discussed in Sec. 3.3.3) and by

the relative ratio of Brønsted acid and basic sites contained within these samples.

Both the intramolecular C=C bond formation (Pathway 2) and isomerization-dehydration

(Pathway 3) reactions involve a single butanal sojourn, during which butanal removes its oxygen

heteroatom by ejecting an H2O molecule while preserving its carbon backbone, according to the

respective chemical equations of:

C3H7CHO+2H C2H5CH=CH2+H2O (Pathway 2 of Scheme 3.1) (3.1)

C3H7CHO CH2=CHCH=CH2+H2O (Pathway 3 of Scheme 3.1) (3.2)

These H2O removal steps are the predominant pathways for alkene and diene formation,

confirmed from the near exclusive formation of CnH2n and CnH2n-2 products from CnH2nO

reactants. The selectivities of CnH2n or CnH2n-2 formation are expressed in terms of the molar

ratios of CnH2n or CnH2n-2 over the total alkene and diene fractions in the product, respectively.

The selectivity values towards CnH2n were 0.95, 0.96, 0.95, and 0.92 for n equals 3, 4, 5, and 6

and towards CnH2n-2 were 0.97, 0.92, and 0.91 for n=4, 5 and 6, respectively (note that n of 3 is

omitted here, because C3H6O reactions do not form C3H4) on H-FAU zeolites at 573 K. In

addition, negligible CO and CO2 were formed under all conditions relevant to deoxygenation

reactions. The carbon selectivities towards CO and CO2, defined by the molar ratio of carbon in

CO and CO2 over the total carbon in the products, are <4% for C3-C6 alkanal reactions on H-

FAU (573 K), <0.03% for butanal reactions on H4SiW12O40 (473-673 K) [6], and <3% for

butanal reactions on H-MFI (573 K) over the entire operating range of our study.

62

The reaction stoichiometry for intramolecular C=C bond formation (Pathway 2 of Scheme 3.1

and Eqn. 3.1) dictates that intermolecular hydrogen transfer, which adds hydrogen atoms to the

alkanals, must occur, before removal of the oxygen atom via dehydration and the eventual alkene

desorption. In the absence of external hydrogen sources, hydrogen atoms are made available

from the cyclization-dehydration-dehydrogenation steps (Steps 1.1.1-1.1.2, 1.3.2-1.3.3, Scheme

3.1). During steady state reaction, H+ sites are occupied by alkanals (in their protonated form) as

the most abundant surface intermediates. This is confirmed from the near stoichiometric butanal-

to-H+ ratios of 1.0, 1.01, and 1.1 on H-MFI (348 K), H-FAU (448 K), and H4SiW12O40 (348 K),

respectively, measured with butanal chemical titrations. Butanal adsorption and H+ site

saturation are also confirmed with Fourier transform infrared spectroscopic studies on H-FAU

zeolites at 348 K; the O-H stretching bands at 3625 cm−1

and 3563 cm−1

, which correspond to the

H+ sites in the supercages and beta cages, respectively, disappear whereas the band at 1675-1685

cm−1

ascribed to the carbonyl stretching band of protonated butanal appears concomitantly [25].

Similarly, the infrared absorption bands of protonated carbonyl group (1670-1690 cm−1

) were

also observed during butanal adsorption on both H4SiW12O40 (348 K) clusters and H-MFI

zeolites (313 K), accompanied by the disappearance of the O-H bands [25]. On such surfaces, the

rates of intramolecular C=C bond formation (4 8,C H OIntrar , per H

+ site), which also equal the site-

time-yields of butene, increase linearly with the total pressure of the aromatic fraction ( AromaticsP ),

as shown in Figure 3.1 for butanal deoxygenation on H-MFI, H-FAU, and H4SiW12O40 at 573 K.

Butanal conversions and partial pressures did not influence the rates for intramolecular C=C

bond formation and their dependences, because these changes on H+ sites that are saturated with

either butanals or their isomers alter neither the identity nor the coverages of the most abundant

surface intermediates during steady-state reactions. This dependence suggests the catalytic

involvement of aromatic species as H-donors. The aromatic fraction contains alkyl benzenes and

alkyl tetralins (e.g., C12H18 and C12H16, respectively), produced from the cyclization-dehydration

(Step 1.3.1) or cyclization-dehydration-dehydrogenation (Step 1.3.2) of alkenal species. These

alkyl benzenes and alkyl tetralins donate their hydrogen atoms and thus further increase their

extents of saturation, forming either alkenyl benzenes or alkyl naphthalenes. Figures 3.2a and

3.2b show the carbon distributions among the aromatic products on H-FAU at different space

velocities, whereas Figures 3.2c and 3.2d depict the carbon distributions on H-MFI and

H4SiW12O40, respectively. The different distributions are results from the different extents in

63

secondary dehydrogenation and transalkylation reactions. For example, dehydrogenation of C8

alkyl cyclohexadienes (C8H12), C13 alkyl benzenes (C13H20), and C16 alkyl tetralins (C16H24)

forms xylenes (C8H10), alkyl naphthalenes (C13H14), and C16 alkyl naphthalenes (C16H20),

respectively. As the space velocity decreases from 0.030 mol butanal (mol H+ s)

–1 to 0.0074 mol

butanal (mol H+ s)

–1, the fractions of C8H10, C13H14, and C16H20 in the aromatic products increase

from 2.6 %, 8.5%, and 8.3 % to 6.3 %, 10.5 %, and 11.8 %, (Figs. 3.2a and 3.2b) because the

lower space velocity and longer contact time favor the secondary dehydrogenation reactions. The

carbon distributions also vary among the different catalysts (Figs. 3.2b-3.2d), because of the

different extents of H+ site confinement, to be discussed in Section 3.3.3.

Figure 3.1. Rates for intramolecular C=C bond formation (Pathway 2, 4 8,C H OIntrar ) as a function of

aromatic pressure ( AromaticsP ) during butanal reactions on H-MFI [■, Si/Al=11.5, space velocity


–1], H-FAU [▲, Si/Al=15, space velocity 0.0074-0.03 mol

butanal (mol H+ s)

−1], and H4SiW12O40 [●, 0.075 mmol H4SiW12O40 gSiO2

−1, space velocity


−1] at 573 K.

0.00 0.02 0.04 0.060

20

40

60

80

H4SiW12O40

H-MFI

H-FAU

r Intr

a, C

4H

8O

(1

0-5

mo

l (m

ol H

+ s

)-1)

PAromatics

(kPa)

64

Figure 3.2. Carbon distributions of aromatic fraction produced in butanal reactions on (a-b) H-

FAU with different space velocities, (c) H-MFI, and (d) H4SiW12O40 at 573 K at time-on-stream

0

5

10

15

C16H

20

C16H

22

C16H

24

C15H

16C

15H

18

C15H

20

C15H

22

C15H

24

C7H

8

C8H

10

C9H

12

C11H

10

C11H

12

C11H

14

C11H

16

C10H

14

C10H

10

C10H

12

C10H

8

Ca

rbo

n d

istr

ibu

tio

n in

aro

ma

tics (

%)

Aromatic products

C12H

12

C12H

14

C12H

16

C12H

18

C13H

14

C13H

16

C13H

18

C13H

20

C14H

22

C14H

20

C14H

18

C14H

16

C17H

22

C17H

24

C17H

26

C18H

24

C18H

26

C18H

28

C19H

30

C19H

28

C19H

26

0

5

10

15

C16H

20

C16H

22

C16H

24

C15H

16

C15H

18

C15H

20

C15H

22

C15H

24

C7H

8

C8H

10

C9H

12

C11H

10

C11H

12

C11H

14

C11H

16

C10H

14

C10H

10

C10H

12

C10H

8

Ca

rbo

n d

istr

ibu

tio

n in

aro

ma

tics (

%)

Aromatic products

C12H

12

C12H

14

C12H

16

C12H

18

C13H

14

C13H

16

C13H

18

C13H

20

C14H

22

C14H

20

C14H

18

C14H

16

C17H

22

C17H

24

C17H

26

C18H

24

C18H

26

C18H

28

C19H

30

C19H

28

C19H

26

0

5

10

15

C16H

20

C16H

22

C16H

24

C15H

16

C15H

18

C15H

20

C15H

22

C15H

24C

7H

8

C8H

10

C9H

12

C11H

10

C11H

12

C11H

14

C11H

16

C10H

14

C10H

10

C10H

12

C10H

8

Ca

rbo

n d

istr

ibu

tio

n in

aro

ma

tics (

%)

Aromatic products

C12H

12

C12H

14

C12H

16

C12H

18

C13H

14

C13H

16

C13H

18

C13H

20

C14H

22 C

14H

20

C14H

18

C14H

16

C17H

22

C17H

24

C17H

26

C18H

24

C18H

26

C18H

28

C19H

30

C19H

28

C19H

26

0

20

40

60

80

C16H

20C

16H

22

C16H

24

C15H

16

C15H

18

C15H

20

C15H

22

C15H

24

C7H

8

C8H

10

C9H

12

C11H

10

C11H

12

C11H

14

C11H

16

C10H

14

C10H

10

C10H

12

C10H

8

Ca

rbo

n d

istr

ibu

tio

n in

aro

ma

tics (

%)

Aromatic products

C12H

12

C12H

14

C12H

16

C12H

18

C13H

14

C13H

16

C13H

18

C13H

20

C14H

22

C14H

20

C14H

18

C14H

16

C17H

22

C17H

24

C17H

26

C18H

24

C18H

26

C18H

28

C19H

30

C19H

28

C19H

26

(c)

(d)

(a)

(b)

H-FAU, space velocity: 0.030 mol butanal (mol H+ s)−1


H-MFI, space velocity: 0.013 mol butanal (mol H+ s)−1

H4SiW12O40, space velocity: 0.18 mol butanal (mol H+ s)−1

65

of 125 min. The distributions include aromatic molecules that do not lose any H ( ) or lose 2

( ), 4 ( ), or 6 ( ) hydrogen atoms in dehydrogenation reactions (e.g., Steps 1.1.2 and 1.3.3,

Scheme 3.1).

The rate dependencies for intramolecular C=C bond formation (4 8,C H OIntrar ) on H

+ sites

predominantly occupied by protonated butanals (Fig. 3.1) are consistent with kinetically relevant

transfer hydrogenation, which shuffles a hydrogen atom from the H-donor (denoted as D) to a

protonated butanal (see derivation and full rate equation in Eqns. S3.1-S3.7, Sec. 3.6.1 of

Appendix). The rate equation, upon simplification, shows that the rate for butanal transfer

hydrogenation, 4 8TH,C H O-Dr , increases linearly with the hydrogen donor pressure, DP :

4 8 4 8TH,C H O- TH,C H O-D D Dr k P (3.3)

where 4 8TH,C H O-Dk is the rate constant for transfer hydrogenation. The linear relations in Figure

3.1 suggest that either aromatics (e.g., methyl- or ethyl-substituted tetralins) or precursors to

aromatics (e.g., 5,6-dimethyl-1,3-cyclohexadiene) act as the H-donors. The transfer

hydrogenation involves cooperative dehydrogenation (Steps 1.1.2 and 1.3.2-1.3.3) of the H-

donors (labeled “D” in Scheme 3.1) and intramolecular C=C bond formation (Steps 2.1-2.2) of

butanal, the H-acceptor (labeled “A” in Scheme 3.1). The reaction leads to alkyl naphthalenes (or

alkyl benzenes) and butene. Only a portion of the aromatic products or precursors of aromatics

can act as H-donors, and depending on their chemical identity, the transfer hydrogenation rate

constant 4 8TH,C H O-Dk varies accordingly. Assuming y is the fraction of a specific H-donor Dy

within the aromatic products (where subscript y denotes the chemical identity) and 4 8TH,C H O- yDk is

the rate constant for transfer hydrogenation between Dy and butanal, the rate for intramolecular

C=C bond formation, 4 8,C H OIntrar , is:

4 84 8 4 8TH,C H O-,C H O ,C H OIntra IntraAromatics Aromatics1 yD

t

yyr k P k P

(3.4a)

4 84 8 TH,C H O-,C H OIntra 1 yD

t

yyk k

(3.4b)

66

According to Equation 3.4a, 4 8,C H OIntrar increases linearly with the total aromatic pressure

( AromaticsP ), consistent with the rate dependency shown in Figure 3.1. 4 8,C H OIntrak is the effective

rate constant for intramolecular C=C bond formation, and it depends on the fraction of H-donors

within the aromatics ( y ) and the transfer hydrogenation rate constants (4 8TH,C H O- yDk ) of the

various H-donors (Dy, y=1,2,…), as shown in Equation 3.4b.

The slopes in Figure 3.1 reflect the rate constants 4 8,C H OIntrak on different catalysts. The rate

constant values were higher on H-FAU than on H-MFI and H4SiW12O40 [4 8,C H OIntrak =13.6±0.3

mmol (mol H+ s kPa)

−1 on H-FAU vs. 5.8±0.3 and 4.3±0.2 mmol (mol H

+ s kPa)

−1 on H-MFI

and H4SiW12O40, respectively, 573 K], indicating that transfer hydrogenation events occur much

more effectively on partially confined, large pore H-FAU zeolites and less so on medium pore H-

MFI zeolites and unconfined structure of H4SiW12O40 clusters. The different reactivities in

transfer hydrogenation among the H-FAU, H-MFI, and H4SiW12O40 catalysts could be caused

either by the difference in H+ site environments among the catalysts or in H-donor identities. In

order to decouple these different contributions, I incorporated either tetralin or cyclohexadiene as

the H-donor into the alkanal reactions and then isolated the rates for tetralin-to-alkanal or

cyclohexadiene-to-alkanal transfer hydrogenation by subtracting the rate contributions of

aromatic products from the total transfer hydrogenation rates, as discussed next in Sections 3.3.2

and 3.3.3.

3.3.2. Mechanism of transfer hydrogenation between tetralins or cyclohexadienes and protonated alkanals at Brønsted acid sites

The transfer hydrogenation between substituted tetralins or cyclohexadienes and protonated

butanals was probed and confirmed by incorporating tetralin, tetralin-adamantane mixture, or

cyclohexadiene in the butanal feed during steady-state butanal deoxygenation reactions. Tetralin

(C10H12) and cyclohexadiene (C6H8, denoted as chd) are known as effective hydrogen donors,

because of their strong thermodynamic tendencies towards dehydrogenation, leading to

naphthalene and benzene, respectively, with more effective π-electron delocalization [26],

whereas adamantane (denoted as ad) is used as a co-catalyst in the transfer hydrogenation

reaction [27]. The rate of each reaction j (j=Inter, Intra, Dehy, or Tish) was measured on H-FAU,

67

H-MFI, and H4SiW12O40 at 573 K while incorporating tetralin (0.08-0.16 kPa), tetralin-

adamantane (0.08-0.16 kPa tetralin and 4-8 Pa adamantane), or cyclohexadiene (0.03-0.15 kPa)

into the butanal feed. These rates with tetralin, tetralin-adamantane, or cyclohexadiene

incorporation (,C H O4 8 -tetralinjr ,

,C H O4 8 -tetralin-adjr , or ,C H O4 8 -chdjr , respectively), when divided by

those in pure butanal feed ( ,C H O4 8jr ), give the rate ratios ,C H O,C H O 4 84 8

1-tetralin

( )j jr r ,

,C H O,C H O 4 84 8

1-tetralin-ad

( )j jr r , or ,C H O,C H O 4 84 8

1-chd

( )j jr r , respectively. These rate ratios are linear

functions of either tetralin pressure ( tetralinP ) or cyclohexadiene pressure ( chd

P ), as shown in

Figure 3.3 for H-FAU zeolites according to:

,C H O

,

,C H O

4 8

4 8

-tetralin

tetralin tetralin1

j

jj

rP

r (3.5a)

,C H O

,

,C H O

4 8

4 8

-tetralin-ad

tetralin-ad tetralin1

j

jj

rP

r (3.5b)

,C H O

,

,C H O

4 8

4 8

-chd

chd chd1

j

jj

rP

r (3.5c)

where ,tetralinj ,

,tetralin-adj , and ,chdj are the proportionality constants and also the slopes of

the data points in Figure 3.3. Their values reflect the extents of promotion for the various

reactions j (j=Inter, Intra, Dehy, or Tish) with tetralin, tetralin-adamantane, or cyclohexadiene

incorporation, as summarized in Table 3.2. For example, the positive ,Intra tetralin

and negative

,Intra tetralin values of 16.1±0.2 and −8.3±0.1 kPa

−1, respectively, on H-FAU zeolites indicate that

tetralin promotes the rate for Pathway 2 (4 8,C H OIntrar ) and inhibits the rate for Pathway 3

(4 8,C H ODehyr ). In contrast, smaller values of

,Inter tetralin (−0.4±0.2 kPa

−1) and

,Tish tetralin

(−1.3±0.3 kPa−1

) on H-FAU correspond to rate changes of less than 10 %; these smaller alpha

values indicate that Pathway 1 (4 8,C H OInterr ) and Pathway 4 ( ,C H O4 8Tishr ) are barely influenced by

tetralin.

68

Figure 3.3. Rate ratios [ ,C H O,C H O 4 84 8

1-tetralin

( )j jr r , ,C H O,C H O 4 84 8

1-tetralin-ad

( )j jr r , or

,C H O,C H O 4 84 8

1-chd

( )j jr r ] for the rates of butanal primary reactions in (a) C4H8O-tetralin

(,C H O4 8 -tetralinjr ), (b) C4H8O-tetralin-adamantane (

,C H O4 8 -tetralin-adjr ), or (c) C4H8O-cyclohexadiene

(,C H O4 8 -chdjr ) feed mixtures to those in C4H8O feed ( ,C H O4 8jr ) as a function of tetralin pressure

( tetralinP ) or cyclohexadiene pressure ( chd

P ) for intermolecular C=C bond formation (Pathway 1,

♦), intramolecular C=C bond formation (Pathway 2, ●), isomerization-dehydration (Pathway 3,

▲), and Tishchenko esterification-ketonization (Pathway 4, ■) on H-FAU at 573 K [subscript

j=Inter, Intra, Dehy, or Tish, which denote inter- or intramolecular C=C bond formation,

isomerization-dehydration, or Tishchenko esterification-ketonization, respectively; space

velocity 0.0074 mol butanal (mol H+ s)

-1, adamantane (if added)=4-8 Pa]. The , j m values

(j=Inter, Intra, Dehy, or Tish; m=tetralin, tetralin-ad, or chd) are determined from the slopes in

these figures by linear regression of the data points against either Equations 3.5a, 3.5b, or 3.5c,

and are summarized in Table 3.2.

0.00 0.02 0.04 0.06 0.080

1

2

3

4

r j,C

4H

8O

-ch

d (r

j, C

4H

8O

)-1

Pchd (kPa)

Intra,chd

0.00 0.05 0.10 0.150

1

2

3

4

r j,C

4H

8O

-te

tra

lin-a

d (r

j, C

4H

8O

)-1 Intra,tetralin-ad

Ptetralin

(kPa)

(b) (c)

0.00 0.05 0.10 0.150

1

2

3

4

r j,C

4H

8O

-tetr

alin

(r j,

C4H

8O

)-1

Intra,tetralin

Ptetralin

(kPa)

(a)

69

Table 3.2. The extent of promotion, ,tetralinj , ,tetralin-adj

, or ,chdj for the various reactions j

(j=Inter, Intra, Dehy, or Tish) with tetralin, tetralin-adamantane, or cyclohexadiene incorporation

during butanal deoxygenation, and the rate constant for cyclohexadiene-to-butanal transfer

hydrogenation, 4 8,C H OTH -chd

k , on H-FAU, H-MFI, and H4SiW12O40 at 573 K

Reaction H-FAU H-MFI H4SiW12O40

Extent of promotion kPa−1

Pathway 1a: 2C4H8O C8H14O+H2O Inter,tetralin

b

Inter,tetralin-ad b

Inter,chd

b

−0.4±0.2

−0.5±0.3

1.2±0.6

1.4±0.1

-

1.2±0.7

0.4±0.2

-

−0.9±0.3

Pathway 2a: C4H8O +2H C4H8+H2O Intra,tetralin

b

Intra,tetralin-ad b

Intra,chd b

16.1±0.2

23.7±0.6

29.0±2.0

1.3±0.2

-

14.4±0.8

0.7±0.3

-

9.2±0.3

Pathway 3a: C4H8O C4H6+H2O Dehy,tetralin

b

Dehy,tetralin-ad

b

Dehy,chd

b

−8.3±0.1

−6.9±0.1

−6.9±0.4

0.1±0.1

-

−28.6±1.0

0±0.3

-

0.6±0.3

Pathway 4a: 2C4H8O C7H14+CO+H2O Tish,tetralin

b

Tish,tetralin-ad b

Tish,chd b

−1.3±0.3

−0.6±0.1

0±0.3

−0.4±0.3

-

−1.1±0.7

1.9±0.4

-

−0.8±0.8

Rate constant mmol (mol H+ s kPa)

−1

C4H8O+C6H8 C4H8+C6H6+H2O , 4 8TH C H O-chdk

c 6.8±0.3 2.8±0.4 0.52±0.03

aPathways 1, 2, 3, and 4 denote the reaction pathways (in Scheme 3.1) of intermolecular C=C bond formation (Step 1),

intramolecular C=C bond formation (Steps 2.1-2.2), isomerization-dehydration (Steps 3.1-3.2), and Tishchenko esterification-

ketonization (Steps 4.1-4.3), respectively;

bThe

, j m values (j= Inter, Intra, Dehy, or Tish; m =tetralin, tetralin-ad, or chd) are the slopes obtained by linear

regression of the data points in Figures 3.3a, 3.3b, or 3.3c against Equations 3.5a, 3.5b, or 3.5c, respectively; cThe rate constants for cyclohexadiene-to-butanal transfer hydrogenation,

,C H O4 8TH -chdk , were measured on H-FAU,

H-MFI, and H4SiW12O40 at 573 K, with space velocities of 0.0074, 0.0033, and 0.045 mol butanal (mol H+ s)

−1,

respectively.

The rate ratio 4 84 8

,C H O,C H O

1Intra Intra-tetralin

( )r r for intramolecular C=C bond formation on H-FAU

zeolites increased with tetralin pressure, thus ,Intra tetralin

acquired a positive value of 16.1±0.2

70

kPa−1

. Incorporation of 4-8 Pa adamantane at an adamantane-to-tetralin ratio of 1:20 increased

the transfer hydrogenation events and the corresponding ,Intra tetralin-ad

value further to 23.7±0.6

kPa−1

, because adamantane acts as a H transfer co-catalyst that donates its tertiary H-atom to

form an adamantyl cation [27, 28], which subsequently abstracts a hydride ion from the tetralin

to complete its catalytic cycle. These additional transfer hydrogenation events as a result of

tetralin addition were confirmed from comparable rates for naphthalene (C10H8) formation (from

tetralin dehydrogenation) and for incremental butene formation. The naphthalene formation rate

is 0.35±0.03 mmol C10H8 (mol H+ s)

−1 with 0.08 kPa tetralin and 1.1 kPa butanal at 573 K on H-

FAU, which corresponds to a H donation rate of 1.4±0.12 mmol H (mol H

+ s)

−1, as each C10H8

turnover donates 4 H atoms (Table 3.3, Steps A and B1). The incremental rate of butene

formation resulted from the tetralin addition, which equals the rate of butanal transfer

hydrogenation by tetralin (4 8,C H OTH -tetralin

r ), was determined by subtracting the rate of butene

formation with C4H8O feed from that with C4H8O-tetralin mixture (4 8,C H OTH -tetralin

r = Intrar =

4 8C H O,Intra -tetralinr −

4 8C H O,Intrar ). The transfer hydrogenation rate, 4 8,C H OTH -tetralin

r , was found to be

0.27±0.02 mmol C4H8 (mol H+ s)

−1 on H-FAU (Table 3.3, Step A), which translates to a H

acceptance rate of 0.54±0.04 mmol H (mol H+ s)

−1, as each C4H8 formation requires 2 H atoms.

This increment in butene formation rates (4 8C H O,Intra -tetralin

r >4 8C H O,Intrar ) and the concomitant

detection of naphthalene upon tetralin addition confirm the involvement of tetralin as the

hydrogen donor. Transfer hydrogenation from alkanes to alkenes has been previously proposed

to involve a hydride ion transfer step at Brønsted acid sites, as established on acidic zeolites

based on density functional theory calculations carried out on a cluster model, H3Si-OH-AlH2-O-

SiH3 [7-9]. The hydride transfer reaction involves an initial alkene protonation that forms a

carbenium ion at the H+ site (CnH2n+H

+CnH2n+1

+). The carbenium ion then accepts a hydride

ion from an alkane (CmH2m+2) via the formation of a carbonium ion (CnH2n+1+∙∙∙H

−∙∙∙CmH2m+1

+)‡

transition state, followed by its decomposition into an alkane (CnH2n+2) and a carbenium ion

(CmH2m+1+). The carbenium ion (CmH2m+1

+) then donates a proton to the zeolitic framework in

order to restore the Brønsted acid site (H+) and desorbs as alkene, CmH2m. I propose a similar

mechanism for the transfer hydrogenation involving tetralin (C10H12) and butanal (C4H8O) as a

hydride donor and acceptor pair at the Brønsted acid site, as shown in Scheme 3.2. Initially, a

butanal adsorbs on a H+ site as protonated C4H8OH

+ (Step I) and, as a result, its carbonyl

71

functional group becomes polarized and the carbonyl carbon acts as a hydride acceptor [29]. The

protonated butanal C4H8OH+

then accepts a hydride ion from tetralin (C10H12) in a concerted step

via a carbonium ion transition state (C4H8OH+∙∙∙H

−∙∙∙C10H11

+)‡ (Step II), which upon

decomposition, leads to a butanol (C4H9OH) molecule and tetralin carbenium ion (C10H11+) (Step

III). The carbenium ion of tetralin (C10H11+) donates a H

+ to the zeolitic framework and forms

1,2-dihydronaphthalene (C10H10), thus regenerating the Brønsted acid site (Step IV) and

completing the catalytic cycle. Butanol then undergoes sequential rapid, kinetically-irrelevant

acid catalyzed dehydration reaction (Step V) that evolves butene (C4H8). The kinetic irrelevance

of the butanol dehydration step (Step 2.2 in Scheme 3.1 and Step V in Scheme 3.2) is confirmed

from the much higher rates for butanol dehydration ( , 4 9Dehy C H OHr ) than for intramolecular C=C

bond formation ( , 4 8C H OIntrar ): the rate ratio for butanol dehydration to the intramolecular C=C

bond formation in butanal, , ,4 9 4 8Dehy C H OH Intra C H O

1( )r r , measured at either 1.1 kPa butanol or butanal,

exceeds 38 on H-MFI and 440 on H4SiW12O40 at 573 K, as summarized in Table 3.1.

Table 3.3. Rates for tetralin dehydrogenation and butanal hydrogenation on H-FAU and

H4SiW12O40 at 573 K

Reactants Butanal (1.1 kPa)-tetralin (0.08 kPa) Tetralin (0.08 kPa) Butanal (1.1 kPa)-H2 (99

kPa)

Reactions Step A: C10H12+2C4H8OC10H8+2C4H8+2H2O

Step B1: C10H12 C10H8+ 2H2

Step B1:

C10H12 C10H8+ 2H2

Step B2:

C4H8O+H2 C4H8+H2O

Rate ,C H O4 8DH -tetralinr a

(Steps A and B1)

,C H O4 8DH -tetralinr b

(Step A)

,DH tetralinr a

(Step B1)

,C H O H4 8 2Hydro -r c

(Step B2)

H-FAU 0.35±0.03 0.27±0.02 0.13±0.01 ~0.007

H4SiW12O40 0.022±0.002 < detection limitd unavailable < detection limit

d

aRate of tetralin (C10H12) dehydrogenation, , DH mr (subscript m denotes either C4H8O-tetralin feed mixture or tetralin), is defined

as the rate of naphthalene (C10H8) formation via Step A and/or Step B1, and given in mmol C10H8 (mol H+ s) −1; space velocity

5.4×10−4 and 3.3×10−3 mol tetralin (mol H+ s) −1 for H-FAU and H4SiW12O40, respectively;

bRate of C4H8O transfer hydrogenation, ,C H O4 8TH -tetralinr , is the rate of C4H8O hydrogenation via Step A in C4H8O-tetralin feed

mixture, and is determined according to Equation 3.8, given in mmol C4H8 (mol H+ s) −1;space velocity 0.0074 and 0.045 mol

butanal (mol H+ s) −1 for H-FAU and H4SiW12O40, respectively;

72

cRate of C4H8O transfer hydrogenation by H2, ,C H O H4 8 2Hydro -r , in C4H8O-H2 feed mixture, is given in mmol C4H8 (mol H+ s) −1;

space velocity 0.0074 and 0.045 mol butanal (mol H+ s) −1 for H-FAU and H4SiW12O40, respectively; dBelow the detection limit.

Butadiene (C4H6) instead of butene may form as a side product from the direct dehydration of

butanal (Pathway 3). The direct dehydration reaction was proposed to occur via allylic alcohol

intermediates [30], as confirmed from the similar yields for the conversions of 2-methylbutanal

and its isomer, methyl isopropyl ketone, to a common intermediate, 2-methyl-2-buten-1-ol,

before its dehydration to isoprene on AlPO4. The yields towards isoprene from 2-methylbutanal

and methyl isopropyl ketone are 49 % and 54 %, respectively, which are very similar. The

primary isomerization-dehydration reaction (Steps 3.1-3.2 in Scheme 3.1 and Eqn. 3.2) converts

butanal to crotyl alcohol before its sequential dehydration to butadiene, according to:

C3H7CHOCH3-CH=CH-CH2OH (also in Step 3.1 of Scheme 3.1) (3.6a)

CH3-CH=CH-CH2OH CH2=CHCH=CH2 + H2O (also in Step 3.2 of Scheme 3.1) (3.6b)

Incorporation of tetralin decreased the net rate of butadiene formation ( ,C H O-4 8Dehy tetralinr <

,C H O4 8Dehyr ), as shown in Figures 3.3a and 3.3b, and led to negative ,Dehy tetralin and ,Dehy tetralin-ad

values (Eqns. 3.5a and 3.5b) of −8.3±0.1 kPa−1

and −6.9±0.1 kPa−1

, respectively. Such decreases

are caused either by tetralin inhibiting the isomerization-dehydration reaction or by the

sequential reaction of butadiene with tetralin. I rule out the former reason because tetralin did not

perturb the concurrent primary pathways (Pathways 1 and 4, Scheme 3.1), an indication that its

addition did not alter the identity of the most abundant surface intermediate. I hypothesize that

butadiene undergoes transfer hydrogenation, during which it accepts two hydrogen atoms from

tetralin and converts to butene (C4H6+C10H12C4H8+C10H10), thus decreasing the net rate of

butadiene formation. In fact, the hydride ion affinity (HIA) for protonated butadiene (C4H7+) is

higher than the carbenium ion of tetralin (C10H11+) (HIA=1011.5 kJ mol

−1 for C4H7

+ vs. 934.1 kJ

mol−1

for C10H11+, Table S3.1, in Sec. 3.6.3 of Appendix), making the transfer hydrogenation

between tetralin and butadiene thermodynamically favourable (to be discussed in Sec. 3.3.3).

Nevertheless, the rate of butadiene (C4H6) transfer hydrogenation by tetralin (4 6,C HTH -tetralin

r ) is

approximately two orders of magnitude lower than the incremental rate of butene formation

73

(, 4 8C H OIntra -tetralin

r − , 4 8C H OIntrar ) [4.8

μmol C4H6 (mol H

+ s)

−1 vs. 270 μmol C4H8 (mol H

+ s)

−1, 573

K, Table 3.4]. Such marked differences in rate magnitude are caused by the much lower

pressures and proton affinity (PA) for butadiene than butanal (1.5 Pa vs. 1.1 kPa and 783.0 kJ

mol−1

vs. 792.7 kJ mol−1

[31], respectively) and by the concomitant lower coverages of

protonated butadiene than butanal at the H+ sites. This result confirms that butadiene

hydrogenation would not affect butanal transfer hydrogenation rate. Similarly, the consumption

of the butene (C4H8) via its transfer hydrogenation that forms butane (C4H10) is also negligible,

as the rate (, 4 8TH C H -tetralin

r ) is less than 0.5 % of the incremental rate of butene formation [1.1

μmol C4H10 (mol H+ s)

−1 vs. 270 μmol C4H8 (mol H

+ s)

−1, Table 3.4].

Table 3.4. Transfer hydrogenation rates of butanal (, 4 8TH C H O-tetralinr ), butadiene (

, 4 6TH C H -tetralinr ), and

butene (, 4 8TH C H -tetralinr ) by tetralin on H-FAU at 573 K

Reaction Rate [mmol (mol H+ s)

−1]

a

C4H8O+C10H12 C4H8+C10H10+H2O , 4 8TH C H O-tetralinr

b

(, 4 8C H OIntra -tetralinr −

,C H O4 8Intrar )

0.27

C4H6+C10H12 C4H8+C10H10 , 4 6TH C H -tetralinr

c

(, 4 8C H ODehyr −

, 4 8C H ODehy -tetralinr )

4.8 ×10−3

C4H8+C10H12 C4H10+C10H10 , 4 8TH C H -tetralinr

d 1.1 ×10

−3

aAll rates were measured on H-FAU at 573 K, 1.1 kPa butanal, 0.08 kPa tetralin (if added), space velocity=0.0074 mol butanal

(mol H+ s) −1, time-on-stream=125 min;

bRate for tetralin-to-butanal transfer hydrogenation, , 4 8TH C H O-tetralinr , is given by the increase in the rate of butene formation upon

tetralin incorporation (, 4 8C H OIntra -tetralinr −

, 4 8C H OIntrar ), according to Equation 3.8;

cRate for tetralin-to-butadiene transfer hydrogenation, , 4 6TH C H -tetralinr , is given by the decrease in the net rate of butadiene

formation upon tetralin incorporation, according to the equation: , 4 6TH C H -tetralinr =

, 4 8C H ODehyr −, 4 8C H ODehy -tetralinr ;

dRate for tetralin-to-butene transfer hydrogenation, , 4 8TH C H -tetralinr , is given by the rate of butane (C4H10) formation.

74

In contrast to the intramolecular C=C bond formation (Eqn. 3.1) and isomerization-dehydration

(Eqn. 3.2) reactions, tetralin incorporation did not perturb the intermolecular C=C bond

formation (Pathway 1 in Scheme 3.1, labeled ♦ in Figs. 3.3a and 3.3b) and Tishchenko

esterification-ketonization (Pathway 4, labeled ■ in Figs. 3.3a and 3.3b) reactions. The extents of

promotion for these reactions, as described by ,tetralinj and

,tetralin-adj (j=Inter or Tish, the

proportionality constants in Eqns. 3.5a and 3.5b), are nearly zero on H-FAU, with or without

adamantane incorporation. ,Inter tetralin

and ,Inter tetralin-ad

are −0.4±0.2 kPa−1

and −0.5±0.3 kPa−1

,

whereas the ,Tish tetralin

and ,Tish tetralin-ad

are −1.3±0.3 kPa−1

and −0.6±0.1 kPa−1

, respectively,

as listed in Table 3.2. The near zero extents of promotion for these reactions (,tetralinj or

,tetralin-adj , j=Inter or Tish) reflect the insensitivity of C8H14O and C7H14 formation rates to

tetralin and adamantane pressures. Taken together, I conclude that tetralin addition and hydride

transfer catalysis alter neither the identity of most abundant surface intermediates nor the

kinetically relevant steps during butanal deoxygenation reactions.

The direct nature of the hydrogen transfer reaction was probed with tetralin, tetralin-butanal, and

hydrogen-butanal reactions on H-FAU (at 573 K) and the hydrogen transfer rates in these

mixtures are summarized in Table 3.3. In the tetralin-butanal mixture (0.08 kPa tetralin, 1.1 kPa

butanal), the rate of tetralin dehydrogenation, defined as the total rate of naphthalene formation

via transfer dehydrogenation and acid catalyzed dehydrogenation (Steps A and B1), is 2.7 times

of that with tetralin (0.08 kPa tetralin) as the sole reactant (Step B1), as shown in Table 3.3. This

result suggests that acid catalyzed tetralin dehydrogenation (Step B1), which occurs via the

initial binding of a H atom from the tetralin to a H+ site in a carbocationic transition state

followed by H2 desorption and the formation of a secondary carbenium ion [32-36], has

comparable but yet slightly smaller reactivities than tetralin transfer dehydrogenation (Step A).

The transfer dehydrogenation involves a direct, intermolecular H transfer step that shuffles

hydride ions (H−) from tetralin to protonated butanal, without the involvement of H2 as the

intermediate. A separate experiment with H2 incorporation showed no rate enhancement for

butene formation: the C4H8O hydrogenation rate with H2 ( , 4 8 2C H O HHydro -r , Step B2, Table 3.3) is

~7 μmol (mol H+ s)

−1. This rate in butanal-H2 mixture is <2% higher than that in butanal (without

H2) and corresponds to a , ,4 8 2 4 8C H O H C H O1

Intra Intra- ( )r r value of 1.02 (90 H2/C4H8O feed ratio, 573

75

K). Thus, diatomic H2, even if formed, did not participate in catalytic turnovers. Based on these

results, I rule out the kinetic significance of the indirect hydrogenation route, which requires an

initial stepwise tetralin dehydrogenation that forms H2 followed by its addition to butanal (Steps

B1 and B2, Table 3.3). Instead, intermolecular H transfer must occur via direct transfer

hydrogenation that shuffles hydride ions from tetralins to protonated butanals.

3.3.3. Catalytic effects of alkanal molecular size and local acid site confinements on transfer hydrogenation reactions

Transfer hydrogenation proposed in Scheme 3.2 involves tetralin (C10H12) and protonated alkanal

(CnH2nOH+) as a hydride donor-acceptor pair in a carbonium ion transition state [7, 9],

(CnH2nOH+∙∙∙H

−∙∙∙C10H11

+)‡. I postulate that the stability of the hydride transfer transition state is

dictated by the difference in hydride ion affinities between the carbenium ions of the hydride

donor (C10H11+) and the hydride acceptor (protonated alkanals, CnH2nOH

+), denoted as

+ +C H C H OH10 11 2- n nHIA :

+ +C H C H OH10 11 2- n nHIA = +C H10 11

HIA − +C H OH2n nHIA (3.7)

+C H10 11HIA and +C H OH2n n

HIA denote the hydride ion affinities of C10H11+ and CnH2nOH

+,

respectively. As an example, the difference in hydride ion affinities between C10H11+ and

C4H8OH+, + +

10 11 4 8C H C H OH-HIA , is −13.7 kJ mol

-1 for the case of butanal deoxygenation. Similarly,

the difference in hydride ion affinities between C10H11+ and either C3H6OH

+, C5H10OH

+, or

C6H12OH+ are –22.5, –10.2, or –7.0 kJ mol

-1 for the cases of propanal, pentanal, and hexanal

deoxygenation, respectively (see Sec. 3.6.3 and Table S3.1 in Appendix for +10 11C H

HIA and

+2C H OHn n

HIA estimation). A more negative + +10 11 2C H C H OH- n n

HIA value indicates a

thermodynamically more stable product and thus a lower activation barrier to evolve the

carbonium ion transition state, as expected from the Brønsted-Evans-Polanyi relation, and a

higher hydride transfer rate. The rates of transfer hydrogenation with tetralin or cyclohexadiene

(chd) as the H-donor and alkanal (CnH2nO, n=3-6) as the H-acceptor, 2,C H OTH -n n Dr (D=tetralin or

chd), are isolated by subtracting the site-time-yields of alkene (CnH2n) in CnH2nO feed (2,C H OIntra n n

r )

76

from those in CnH2nO-tetralin or CnH2nO-cyclohexadiene feed mixture (2,C H OIntra -n n Dr ), measured

on H-FAU zeolites at 573 K:

22 2 ,C H O,C H O ,C H OTH Intra Intra- - nn n nn nD Dr r r (3.8)

2,C H OTH -n n Dr is proportional to the H-donor pressure DP (D=tetralin or chd) via a proportionality

constant 2,C H OTH -n n Dk , which is also the rate constant for transfer hydrogenation between the H-

donor D and alkanals CnH2nO:

,C H O2TH -n n Dr = 2,C H OTH -n n D Dk P (3.9)

as also indicated from the linear relation in Equations. 3.5a-3.5c and shown in Figures 3.3a-3.3c.

2,C H OTH -n n Dk is also the rate constant for the kinetically relevant nucleophilic attack of the H-

donor D (e.g. tetralin) onto the protonated alkanal (e.g. C4H8O) via the formation of bimolecular

carbocationic transition state, as shown in Step II of Scheme 3.2.

77

Scheme 3.2. A proposed mechanism for intermolecular hydride transfer from tetralin to

protonated butanal (the kinetically relevant step for Pathway 2).

It was noted that the reactions of alkanal (CnH2nO)-tetralin (C10H12) mixtures on H-FAU

produced alkyl tetralins (Cn+10H2n+12) and alkyl naphthalenes (Cn+10H2n+8), which, however, were

not detected in the reaction of alkanal (CnH2nO) alone. It is known that alkene can undergo

electrophilic alkylation reaction with aromatics on acidic zeolites [37-41]. Therefore, I speculate

that alkene (CnH2n) formed from alkanal transfer hydrogenation could undergo alkylation

reaction with either the tetralin (C10H12) or its dehydrogenation product, naphthalene (C10H8), to

produce alkyl tetralins (Cn+10H2n+12) or alkyl naphthalenes (Cn+10H2n+8), respectively, as shown in

Steps AK1 and AK2 in Scheme 3.3. These alkylation reactions consumed the alkene (CnH2n) and

decreased the net rate of alkene formation from alkanal transfer hydrogenation. Therefore, the

rate for the intramolecular C=C bond formation in the alkanal-tetralin feed mixture,

78

2,C H OIntra -tetralinn nr , was determined from the total rates of alkene (CnH2n), alkyl tetralin

(Cn+10H2n+12), and alkyl naphthalene (Cn+10H2n+8) formation.

Scheme 3.3. Reaction network for CnH2n-tetralin (naphthalene) alkylation.

Figure 3.4 shows the rate constant 2,C H OTH -tetralinn n

k for different alkanals (CnH2nO, n=3-6) as a

function of the hydride ion affinity differences, + +10 11 2C H C H OH- n n

HIA . As the carbon number of

alkanal decreases from six to four, their hydride ion affinity ( +2C H OHn n

HIA ) increases from 941.1

to 947.8 kJ mol−1

and thus the + +10 11 2C H C H OH- n n

HIA concomitantly decreases from –7.0 to –13.7 kJ

mol−1

(Table S3.1, in Sec. 3.6.3 of Appendix). The more negative hydride ion affinity difference

leads to stronger CnH2nOH+∙∙∙H

− bond and more stable (CnH2nOH

+∙∙∙H

−∙∙∙C10H11

+)‡ structure at

the transition state and in turn to higher 2,C H OTH -tetralinn n

k values. This trend, however, does not

apply for propanal (C3H6O). Despite the more negative hydride ion affinity difference for

propanal than butanal ( + +10 11 3 6C H C H OH-

HIA = −22.5 kJ mol−1

vs. + +10 11 4 8C H C H OH-

HIA = −13.7 kJ

mol−1

), the rate of transfer hydrogenation is lower [3 6,C H OTH -tetralin

k =2.21±0.03 mmol (mol H+ s

kPa)−1

vs. 4 8,C H OTH -tetralin

k =4.3±0.2 mmol (mol H+ s kPa)

−1, Fig. 3.4]. This exceptional activity

79

trend for propanal transfer hydrogenation is probably related to the solvation effects of the H+

site local confinement on the stability of the hydride transfer transition state.

Figure 3.4. Rate constants for tetralin-to-alkanal (CnH2nO) transfer hydrogenation

(2,C H OTH -tetralinn n

k , Eqn. 3.9, n=3-6) on H-FAU as a function of the hydride ion affinity difference

between the protonated alkanal (CnH2nOH+) and the carbenium ion of tetralin (C10H11

+) (the

hydride ion affinity difference is given by Eqn. 3.7) [573 K, 1,1 kPa alkanal, 0.08-0.16 kPa

tetralin, space velocity 0.0074 mol alkanal (mol H+ s)

−1].

Kinetic investigations of hydride transfer from isobutane to C3 and C6 alkoxides during propene

oligomerization on solid Brønsted acids show a direct relation between hydride transfer rates and

the stabilities and concentrations of hydride ion acceptors [42]. The hydride transfer to C3 and C6

alkoxides that forms propane and hexane, respectively, is much less effective for C3 than C6

alkoxides, because of the smaller van der Waals stabilizations and thus lower concentrations of

C3 than C6 alkoxides [42]. These hydride transfer reactivity trends for C3 and C6 alkoxides [42]

are in the exact opposite to those between alkanal and tetralin (in Fig. 3.4), because their

kinetically relevant steps differ: the hydride transfer between isobutane and C3 (or C6) alkoxides

during propene oligomerizations is governed by the formation of alkoxides as the H-acceptors

-25 -20 -15 -10 -5 00

1

2

3

4

5

6

kT

H,C

nH

2nO

-te

tra

lin(m

mol (m

ol H

+ s

kP

a)-1

)

HIAC10H11+CnH2nOH+ (kJ mol

-1)

n=6

n=5

n=3

n=4

80

[42], but this is not the case between tetralin and C3-C6 alkanals. The latter is not limited by the

formation of H-acceptor (CnH2nOH+), because the H

+ sites are already populated by CnH2nOH

+

as the most abundant surface intermediates. Instead, it is limited by the bimolecular reactions

between the hydride donor and acceptor pair to evolve the carbonium ion transition state

(CnH2nOH+∙∙∙H

−∙∙∙C10H11

+)‡ (Step II, Scheme 3.2). The preferential occupation of the H

+ site by

alkanals is confirmed from chemical titrations with alkanals [1, 6], first order dependence of the

parallel, bimolecular alkanal condensation reaction on the alkanal pressure (Step 1, Scheme 3.1)

[1, 6], and the infrared spectroscopic study of butanal adsorption on H-FAU, which shows the

disappearance of the O-H stretching bands at 3625 cm−1

(in supercages) and 3563 cm−1

(in beta

cages) and the concomitant appearance of C=O stretching bands at 1675-1685 cm−1

[25].

One could not rule out the effects of local structural confinements on these reactivity trends

within the alkanal homologues, as such confinements may stabilize the alkanals as well as the

transition states: it is plausible that (C4H8OH+∙∙∙H

−∙∙∙C10H11

+)‡ transition state may be stabilized

more than (C3H6OH+∙∙∙H

−∙∙∙C10H11

+)‡ through larger van der Waals interactions, thus leading to

the higher transfer hydrogenation reactivities for C4H8O than C3H6O (Fig. 3.4). Similar solvation

effects caused by local confinements have been reported for hydride transfer in solid acid-

catalyzed dimethyl ether homologation [14] and n-pentane cracking [13]. In dimethyl ether

homologation, large-pore FAU and BEA zeolites exhibit higher selectivities to triptane than

mesoporous SiO2-Al2O3 and medium-pore MFI zeolites, because the pores and cages in FAU

and BEA zeolites provide the required physical dimensions as the “right fit” that solvate and

stabilize the hydride transfer and methylation transition states to a larger extent than the

mesoporous SiO2-Al2O3 and medium-pore MFI zeolites, thus favoring the chain termination to

C7 product (triptane) [14]. In n-pentane cracking, zeolites with cavity volumes of about 0.2 nm3

[e.g., CIT-1 (0.211 nm3) and MCM-68 (0.182 nm

3)] were found to be the most active ones in

catalyzing the hydride transfer from n-C5H12 to C2H5+, C3H7

+, or C4H9

+carbenium ions, which

forms alkanes (e.g., C2H6, C3H8, or C4H10, respectively), likely because the cavity volumes of

these zeolites and the bimolecular transition states are of comparable dimensions: the volumes

for C2H5+-n-C5H10, C3H7

+-n-C5H10, and C4H9

+-n-C5H10 fragments were estimated to be 0.182,

0.203, and 0.224 nm3, respectively [13]. These effects of site confinement on alkanal transfer

hydrogenation are probed with butanal reactions on H+ sites confined to different extents -in

medium pore straight or zig-zag channels of MFI structures (5.4 Å), in supercages (diameter of

81

11.8 Å) and 12-membered ring windows (7.4 Å) of FAU structures, or dispersed on the surfaces

of H4SiW12O40 clusters that are then immobilized on mesoporous silica (without molecular scale

confinements). Partially confined FAU structures exhibit much higher transfer hydrogenation

reactivities than the smaller pore MFI zeolites and unconfined H4SiW12O40 clusters. First, H-

FAU promotes tetralin-to-butanal transfer hydrogenation to the largest extent among these

catalysts, as indicated by the much larger ,Intra tetralin

values (Eqn. 3.5a, 16.1±0.2 kPa−1

on H-

FAU vs. 1.3±0.2 kPa−1

and 0.7±0.3 kPa−1

on H-MFI and H4SiW12O40, respectively, Table 3.2).

Second, the selectivities to intramolecular C=C bond formation (Pathway 2), 4 8,C H OIntraS , are

much higher on H-MFI and H-FAU than on H4SiW12O40 (4 8,C H OIntraS =0.27 and 0.37 vs. 0.01,

Table 3.1), an indication that butanal transfer hydrogenation is promoted within the confined

pores and cages of molecular dimensions than on the unconfined H4SiW12O40 clusters. I

hypothesize that the partially confined environment in H-FAU (diameter of supercages=11.8 Å)

could stabilize the (C4H8OH+∙∙∙H

−∙∙∙C10H11

+)‡ carbonium ion type transition state to a greater

extent than H-MFI and H4SiW12O40. The pores of H-MFI (5.4 Å) are too small to accommodate

the bimolecular transition state complex consisting of fragments from a large H-donor (e.g.,

tetralin, kinetic diameter >6.2 Å [43]) and the protonated butanal (~5.0 Å). Such a structure

permits only smaller compounds acting as the H-donors (e.g., 5,6-dimethyl-1,3-cyclohexadiene),

thus leading to lower transfer hydrogenation reactivities and lower selectivities on H-MFI than

on H-FAU (4 8,C H OIntraS =0.27 on H-MFI vs. 0.37 on H-FAU, Table 3.1).

The fraction of C7-C9 aromatics formed upon the dehydrogenation-transalkylation of C8

cycloalkadienes, is higher on H-MFI (16 %, Fig. 3.2c) than on H-FAU (10 %, Fig. 3.2a). In

contrast, the fraction of the dehydrogenated C11-C17 aromatic species is much lower on H-MFI

(45 %, Fig. 3.2c) than on H-FAU (70 %, Fig. 3.2a). These distinct product distributions suggest

the difference in molecular dimension and identity of the hydrogen donors: the primary H-donors

in medium pore H-MFI are the smaller hydrocarbon species (e.g. C8 alkyl cyclohexadienes),

whereas in the large pore H-FAU are larger aromatics, ranging from C8 to C17.

This size requirement for H-donor was probed by incorporating cyclohexadiene or tetralin as the

H-donor for butanal transfer hydrogenation. The kinetic diameter of cyclohexadiene is estimated

to be ~5.9 Å [44], therefore it can access both the medium pores (5.4 Å) of H-MFI and the large

82

pores (11.8 Å) of H-FAU and act as an effective H-donor on both catalysts ( ,chdIntra =14.4±0.8

and 29.0±2.0 kPa−1

on H-MFI and H-FAU, respectively, Table 3.2). In contrast, tetralin (~6.2 Å

[43]) can access only the large pores of H-FAU but not the medium pore of H-MFI; as a result,

tetralin promotes only the butanal transfer hydrogenation on H-FAU ( ,Intra tetralin =16.1±0.2 and

1.3±0.2 kPa−1

on H-FAU and H-MFI, respectively, Table 3.2). The rate constant for

cyclohexadiene-to-butanal transfer hydrogenation, 4 8TH,C H O-chdk , is lower on H-MFI than on H-

FAU [2.8±0.4 vs. 6.8±0.3 mmol (mol H+ s kPa)

−1, Table 3.2]. This might be caused by the steric

constraints in H-MFI that destabilize the bulky transition state (C4H8OH+∙∙∙H

−∙∙∙C6H7

+)‡ of

hydride transfer, as the pores of H-MFI (5.4 Å) have similar dimensions with the kinetic

diameter of cyclohexadiene (~5.9 Å). The steric constraints may partially compensate the

solvation effects, making the transition state less stable in the medium pores of H-MFI than the

supercages of H-FAU.

In contrast to the H+ sites confined within the microporous structure of H-MFI and H-FAU

zeolites, the unconfined H+ sites on H4SiW12O40 clusters produce strictly aromatics with 12 or 16

carbon atoms (Fig. 3.2d) as a result of cyclization-dehydration of trimeric (C12H20O) and

tetrameric (C16H26O) alkenals produced from the intermolecular C=C bond formation steps,

respectively. These aromatics do not undergo sequential transalkylation on H4SiW12O40 clusters,

as transalkylation requires confined reaction environment. The lack of confinement in

H4SiW12O40 clusters also inhibits the hydride transfer step, thus only a small fraction (22 %) of

the aromatics undergo dehydrogenation, as shown in Figure 3.2d. For this reason, tetralin does

not promote butanal transfer hydrogenation on H4SiW12O40, resulting in a small ,Intra tetralin value

of 0.7±0.3 kPa−1

(Table 3.2). Cyclohexadiene as a more effective H-donor promotes butanal

transfer hydrogenation on H4SiW12O40, resulting in a positive ,Intra chd value of 9.3±0.3 kPa−1

(Table 3.2), because its carbenium ion has a much lower hydride ion affinity than that of tetralin

( +76C H

HIA = 907.2 kJ mol−1

vs. +10 11C H

HIA = 934.1 kJ mol−1

, Table S3.1, Appendix); thus, hydride

abstraction is energetically more favorable from cyclohexadiene than tetralin. Despite the

promotional effects, the rate constant for cyclohexadiene-to-butanal transfer hydrogenation,

4 8TH,C H O-chdk , is much lower on H4SiW12O40 than on H-FAU and H-MFI [0.52±0.03 vs. 6.8±0.3

83

and 2.8±0.4 mmol (mol H+ s kPa)

−1, respectively, Table 3.2], an indication that the unconfined

H+ sites on H4SiW12O40 are much less effective for transfer hydrogenation.

3.4. Conclusion

The transfer hydrogenation of alkanal (CnH2nO) occurs on Brønsted acid sites via a kinetically

relevant hydride transfer step that shuffles a hydride ion from a hydride donor to a protonated

alkanal, followed by rapid H2O removal and the concomitant creation of an intramolecular C=C

bond in the alkanal reactant. The reaction forms an alkene (CnH2n) in catalytic sojourns that

retain the carbon backbone of alkanal. The hydride transfer step involves a bimolecular

(CnH2nOH+∙∙∙H

−∙∙∙R’H

+)‡ transition state consisted of a hydride donor (hydrocarbon species, R’H2)

and acceptor (protonated alkanal, CnH2nOH+) pair, with both fragments sharing a hydride ion,

similar to those previously found for the hydride transfer between alkane and alkoxide (or the

related carbenium ion) in confined, microporous structures. The rates of hydride ion transfer vary

linearly with the partial pressure of H-donors, identified to be the larger hydrocarbon species

(e.g., alkyl tetralins), produced from the parallel intermolecular C=C bond formation and ring

closure steps, and confirmed here by tetralin, cyclohexadiene, and adamantane (a co-catalyst)

incorporation which specifically promotes transfer hydrogenation among the various pathways of

alkanal deoxygenation.

The rate constants for transfer hydrogenation are correlated to the hydride ion affinity difference

between the carbenium ions of the H-donors (R’H+) and the protonated alkanals (CnH2nOH

+), as

it dictates the heat of the hydride transfer reaction and in turn the stability of the hydride transfer

transition state. Smaller alkanals with higher hydride ion affinities exhibit larger transfer

hydrogenation rates (C4>C5>C6). An exception is propanal-despite its higher hydride ion affinity,

it appears to be much less reactive in the transfer hydrogenation than butanal, a phenomenon

likely caused by the smaller transition state for propanal than for larger alkanals and its lesser

extent of solvation in the FAU cages.

The transfer hydrogenation occurs much faster on H+ sites confined within the supercages of

FAU zeolites than those in the medium pore MFI zeolites or dispersed on unconfined

H4SiW12O40 Keggin structures, likely because partially confined environments in the FAU

84

supercages solvate and stabilize the bimolecular (CnH2nOH+∙∙∙H

−∙∙∙R’H

+)‡ transition state via van

der Waals interactions. As the alkanal size increases from C3 to C6, the transition state may be

stabilized by their larger van der Waals interactions with the zeolitic wall, but such stabilizations

are offset by the destabilization caused by steric hindrances. I hypothesize that the interaction

between the transition state and local site environment, together with the variation in the hydride

ion affinity of alkanals, leads to the observed reactivity trend of C4>C3>C5>C6 for tetralin-to-

alkanal transfer hydrogenation in H-FAU zeolites. These effects of local site structures and the

thermochemical properties of reactant determine the reactivity of hydride transfer within the

various concomitant catalytic cycles and thus govern the selectivity ratio towards alkenes, dienes,

aromatics, and larger oxygenates during alkanal deoxygenation reactions.

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[23] H. Kawakami, S. Yoshida, Quantum chemical studies of alumina. Part 3.-Acidity and

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Sites on Polyoxometalate Clusters, Chem. Cat. Chem., to be submitted for publication.

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88

3.6. Appendix

3.6.1. Rate equation for intramolecular C=C bond formation

Butanal and all the other species (denoted as x) are adsorbed at the H+ sites and their surface

concentrations are 4 8[C H O*] and [Other*] , respectively. These surface concentrations are

derived from the respective quasi-equilibrated adsorption step for each molecule:

4 84 8

+C H Oads,C H O4 8

[C H O*] [H ]K P (S3.1)

+ + +ads, ads,1 1

[Other*] [H ] [H ] [H ]n n

x xx xx xK P K P S

(S3.2)

ads,1

n

xxxS K P

(S3.3)

where 4 8ads,C H O

K and ads,xK are the equilibrium constants for C4H8O and species x adsorption

at H+ sites, respectively; [H

+] is the concentration of free H

+ sites;

4 8C H OP and xP are the partial

pressures of C4H8O and species x, respectively. The total H+ site concentration, [T], is given by:

+4 8

[T] [H ] [C H O*] [Other*] (S3.4)

The transfer hydrogenation of C4H8O occurs when an adsorbed C4H8O and a H-donor (denoted

as D) from the gas phase reacts in a kinetically relevant step (Step II, Scheme 3.2). The rate

,C H O4 8TH -Dr is:

,C H O ,C H O ,C H O4 8 4 8 4 8TH TH TH- - - +

4 8 4 8

4 8

[C H O*] [C H O*]

[T] [H ] [C H O*] [Other*]D DD D Dr k P k P

(S3.5)

where ,C H O4 8TH -Dk is the rate constant for the D-to-C4H8O transfer hydrogenation, and DP is the

partial pressure of the H-donor. By substituting Equations S3.1 and S3.2 into S3.5, the rate

equation becomes:

89

,C H O

,C H O

4 8

4 8

TH

TH4 84 8

4 84 8

- C H Oads,C H O

-C H Oads,C H O

1

DD

D

k K P Pr

K P S

(S3.6)

When C4H8O is the most abundant surface species, the term 4 84 8 C H Oads,C H O

K P is much larger

than the other terms in the denominator [4 84 8 C H Oads,C H O

K P >> (1+S)], thus Equation S3.6 is

simplified to Equation S3.7:

,C H O ,C H O4 8 4 8TH TH- - DD Dr k P (S3.7)

which is also Equation 3.3 in Sec. 3.3.1.

3.6.2. Intermolecular C=C bond formation and Tishchenko esterification of butanal on NaH-MFI zeolites

Tishchenko esterification (Step 4.1, the primary step of Pathway 4) of aldehydes is known to

occur on solid base catalysts (e.g. benzaldehyde [45] and butanal [5, 46] on alkaline earth oxides

[5, 45, 46] and γ-Al2O3 [46]). The active site requirements for the pathways of Tishchenko

esterification (Pathway 4, Scheme 3.1) and intermolecular C=C bond formation (Pathway 1,

Scheme 3.1) were examined using MFI zeolites with different acidic and basic site ratios. A

series of MFI zeolites (Si/Al=11.5) with different amounts of acidic and basic sites were

synthesized by Na-exchange, and their acid site densities were measured by pyridine titration at

473 K. The quantities of Na+ sites on these samples were determined by subtracting the

remaining acid sites from the original acid site density (1.12 mmol gcat.−1

), while the amount of

extra-framework alumina on this MFI sample was measured to be 0.18 mmol gcat.−1

via infrared

spectroscopic study of pyridine adsorption, based on the band of coordinated pyridine at 1455

cm−1

, as reported in our previous work [1]. In these samples, the conjugated oxygen of Na+ site

[24] and the bi-coordinated oxygen of the extra-framework alumina [23] may both act as the

basic sites.

Figure S3.1a shows the rates for intermolecular C=C bond formation ( ,C H O4 8Interr ) and

Tishchenko esterification-ketonization ( ,C H O4 8Tishr ) on these series of samples (573 K) as a

function of H+ site density. ,C H O4 8Interr increases proportionally with H

+ site concentration,

90

despite the decrease in the basic site (Na+) density. This result indicates that the intermolecular

C=C bond formation (Pathway 1) occurs strictly at the H+ sites. In contrast, ,C H O4 8Tish

r does not

increase with the H+ site concentration but instead it increases proportionally with the basic site

concentration (Fig. S3.1b). These site dependence results confirm that Tishchenko esterification-

ketonization (Pathway 4) is catalyzed strictly by the basic sites, at least in the experimental

conditions reported in our study, and consistent with those established in the literature [5, 45, 46].

Figure S3.1. (a) The rates for intermolecular C=C bond formation ( ,C H O4 8Interr , ♦) and

Tishchenko esterification-ketonization ( ,C H O4 8Tishr , □) as a function of H

+ site density and (b) the

rates for Tishchenko esterification-ketonization ( ,C H O4 8Tishr , □) as a function of basic site density

during butanal reaction on H-MFI and NaH-MFI zeolites at 573 K [1.1 kPa butanal, space

velocity=0.0037 mmol butanal (gcat. s) −1

, time-on-stream=125 min].

3.6.3. Estimation of hydride ion affinities for protonated alkanals and carbenium ions of H-donors

The hydride ion affinities (HIA) for the carbenium ions (R’H+) of the selected hydrocarbons

(R’H2, e.g., tetralin C10H12) and the protonated alkanals (CnH2nOH+, n=3, 4, 5, 6) (denoted as

0.0 0.5 1.0 1.50

1

2

3

4

5

rTish,C4H8O

Basic site density (10-3

mol gcat.-1

)

rInter,C4H8O

Ra

te (

10

-8 m

ol (g

cat. s

)-1)

H+ site density (10

-3 mol gcat.

-1)

1.0 0.5 0.0

0.0 0.2 0.4 0.6 0.8 1.0 1.20

1

2

3

4

5

rTish,C4H8O

Ra

te (

10

-8 m

ol (g

cat. s

)-1)

Basic site density (10-3

mol gcat.-1

)

(a) (b)

91

+R'HHIA and +

2C H OHn nHIA , respectively) are defined as the heats of hydride ion addition

reactions as shown in Equations S3.8 and S3.9, respectively:

R’H++ H

−→RH2 (S3.8)

CnH2nOH++ H

−→CnH2n+1OH (S3.9)

+R'HHIA and +

2C H OHn nHIA were estimated based on the thermochemical cycles shown in

Scheme S3.1, which use tetralin (Scheme S3.1a) and alkanal (Scheme S3.1b) to illustrate the

method. The values of +R'HHIA (or +

2C H OHn nHIA ) were calculated according to Equations S3.10-

S3.13, using proton affinities (PA), heats of formation ( fH ) of the hydrogen donors R’H2 (or

the alkanal hydrogenation products CnH2n+1OH), the diatomic hydrogen H2, and the

dehydrogenation products R’ (or the alkanals, CnH2nO), and heat of reaction for the reaction

between a proton and a hydride (H++H

−→H2, r H ion

H =−1675.3 kJ mol-1

[47]). The estimated

values are listed in Table S3.1.

92


the carbenium ion (R’H+) of a hydrocarbon (R’H2, taking R’H2=tetralin as an example) and (b)

the protonated alkanal (CnH2nOH+) [ R'PA and

2C H On nPA are the proton affinities of hydrocarbon

R’ and alkanal CnH2nO, respectively; r H ionH (−1675.3 kJ mol

-1) [47] is the heat of reaction

for H++H

− → H2; r Hydro,R'

H and 2

r Hydro,C H On nH are the heats of reaction for hydrogenation

reactions: R’+H2→ R’H2 and CnH2nO+ H2→CnH2n+1OH, respectively].

+ r rR' H ion Hydro,R'R'HHIA PA H H (S3.10)

and

22r HR'H R'Hydro,R' f f fH H H H (S3.11)

+2 22

r rC H O H ion Hydro,C H OC H OH n nnn nnHIA PA H H (S3.12)

and

22 1 22r HC H OH C H OHydro,C H O f f fn nn n nnH H H H

(S3.13)

(R’H2)

(R’)

(R’H+)

(R’)

(CnH2n+1OH)(CnH2nOH+)

(CnH2nO) (CnH2nO)

(a) (b)

93

Table S3.1. Hydride ion affinities (HIA) for protonated alkanals (CnH2nOH+; n=3-6) and the

carbenium ions of the H-donors (R’H+)

Carbenium ion

(R’H+) or protonated

alkanal (CnH2nOH+)

H-donor (R’H2) or

alkanal (CnH2nO)

R'PA or

2C H On nPA

(kJ mol−1

)a

R'f H or

2C H Of n nH

(kJ mol−1

)b

2R'Hf H or

2 1C H OHf n nH

(kJ mol−1

)b

+R'HHIA or

+2C H OHn n

HIA

(kJ mol−1

)

C3H6OH+ Propanal 786.0 -188.7 -256 956.6

C4H8OH+ Butanal 792.7 -211.8 -277 947.8

C5H10OH+ Pentanal 796.6 -232.4 -298 944.3

C6H12OH+ Hexanal 801.6 -248.6 -316 941.1

C10H11+ Tetralin 842.0 [48] 130.8 30.0 934.1

C6H7+ Cyclohexadiene 746.4 82.9 104.6 907.2

C3H7+ Propane 751.6 20.41 -104.7 1048.8

C4H9+ Butane 802.1 -17.9 -134.2 989.5

C5H11+ Pentane 808.8 -41.5 -153.7 978.7

C6H13+ Hexane 812.9 -63.51 -171.6 970.5

C4H7+ Butene 783.4 108.8 -10.8 1011.5

aProton affinities (PA) were obtained from ref. [31, 47]. bHeats of formation were obtained from the database [47].

3.6.4. Effects of time-on-stream on the amount of remaining H+

sites, the rate of butanal conversion, and the reaction selectivities on H-MFI, H-FAU, and H4SiW12O40 catalysts

Figures S3.2a, S3.2b, and S3.2c show the remaining amount of H+ sites on H-MFI, H-FAU, and

H4SiW12O40, respectively, after butanal reactions (573 K) at different time-on-streams, measured

by chemical titration with pyridine at 473 K. Reactivities decreased after butanal reactions, due

to the formation of heavier products (e.g. larger aromatics and coke) which occupied the H+ sites.

On H-MFI, the H+ site density decreased from 1.12 to 0.75 mmol H

+ gcat.

−1 (Fig. S3.2a) during

155 min of butanal reaction. On H-FAU, the amount of acid sites decreased from 0.506 to 0.214

mmol H+ gcat.

−1, during the initial 125 min, but the deactivation was much slower after 125 min

94

(Fig. S3.2b). H4SiW12O40 deactivated quickly within the initial 35 min but stayed relatively

stable afterward, with H+ site density maintained at 0.095±0.017 mmol H

+ gcat.

−1 (Fig. S3.2c).

Figures S3.3a, S3.3b, and S3.3c show butanal conversion rates (4 8,C H Ooverallr ) and carbon

selectivities to different reaction pathways (4 8,C H OjS , j=Inter, Intra, Dehy, or Tish) on H-MFI, H-

FAU, and H4SiW12O40, respectively, as a function of time-on-stream. Here 4 8,C H OjS is defined as

the rate of butanal consumption in reaction j divided by the overall butanal conversion rate

4 8,C H Ooverallr . During the initial 125 min, the butanal conversion rates (4 8,C H Ooverallr ) on H-MFI and

H-FAU decreased by >72% and >47%, respectively, and the carbon selectivities (4 8,C H OjS )

commensurately changed (as shown in Figs. S3.3a and S3.3b), because of the loss of H+ site

(Figs. S3.2a and S3.2b) caused by the formation of heavier products (e.g. larger aromatics and

coke) inside the zeolitic pores. The changes in rate per unit time, defined as 4 8,C H Ooverallr (∆time-

on-stream) −1

, were one order of magnitude slower than the initial values for time-on-stream

above 125 min and 4 8,C H OjS were less than ±6% during 240 min of measurement for time-on-

stream above 125 min (Figs. S3.3a and S3.3b). On H4SiW12O40, 4 8,C H Ooverallr reached steady-state

after 35 min, and the carbon selectivities became stable above 125 min (4 8,C H OjS <5 %), as

shown in Fig. S3.3c.

Figure S3.2. Concentration of remaining H+ sites on (a) H-MFI, (b) H-FAU, and (c) H4SiW12O40

catalysts after exposure to butanal reactants for different reaction times [573 K, space

0 50 100 150 200 2500.0

0.2

0.4

0.6

0.8

Rem

ain

ing

H+ s

ites (

103m

ol g

cat.1)

Time-on-stream (min)

0 50 100 150 200 2500.0

0.5

1.0

1.5

Rem

ain

ing

H+ s

ites (

103m

ol g

cat.1)


0 100 200 700 8000.00

0.05

0.10

0.15

0.20

Re

ma

inin

g H

+ s

ite

s (

103m

ol g

cat.1)


(a) (b) (c)

95

velocity=0.0037-0.015, 0.0037, and 0.0074-0.030 mmol butanal (gcat. s)-1

for H-MFI, H-FAU,

and H4SiW12O40, respectively].

Figure S3.3. Butanal conversion rates (4 8,C H Ooverallr , ◊) and selectivities to intermolecular C=C

bond formation (4 8,C H OInterS , ●), intramolecular C=C bond formation (

4 8,C H OIntraS , ▲),

isomerization-dehydration (4 8,C H ODehyS , ▼), and Tishchenko esterification-ketonization (

4 8,C H OTishS ,

■) during butanal reactions on (a) H-MFI, (b) H-FAU, and (c) H4SiW12O40 at 573 K as a

function of time-on-stream [1.1 kPa butanal, space velocity=0.0033, 0.0074, and 0.045 mol

butanal (mol H+ s)

−1 for H-MFI, H-FAU, and H4SiW12O40, respectively].

3.6.5. Parity plots for the kinetic data

Figure S3.4 shows the parity plots for the predicted and measured rates (Figure 3.1) for the

intramolecular C=C bond formation, 4 8,C H OIntrar , during butanal reaction at 573 K, on H-MFI, H-

FAU, and H4SiW12O40, respectively. The predicted 4 8,C H OIntrar were determined via Equation

S3.14:

4 8 4 8,C H O ,C H OIntra Intra Aromaticsr k P (S3.14)

which is also Equation 3.4a in Section 3.3.1. Here the rate constant, 4 8,C H OIntrak , was determined

by the linear regression of the kinetic data in Figure 3.1 against Equation 3.4a in Section 3.3.1.

0 100 200 300 4000.0

0.5

1.0

1.5

STish,C4H8O

SDehy,C4H8O

SIntra,C4H8O

SInter,C4H8O

r overa

ll,C

4H

8O

(1

06m

ol (g

cat. s

)1)


roverall,C4H8O

0

20

40

60

Ca

rbo

n s

ele

ctivitie

s (

%)

0 200 400 600 800 10000.0

0.2

0.4

0.6

0.8

1.0

1.2

STish,C4H8O

SDehy,C4H8O

SIntra,C4H8O

SInter,C4H8O

r overa

ll,C

4H

8O

(1

06m

ol (g

cat. s

)1)


roverall,C4H8O

0

20

40

60

Carb

on

sele

ctivitie

s (

%)

0 100 200 3000.00

0.05

0.10

0.15

0.20

STish,C4H8O

SDehy,C4H8O

SIntra,C4H8O

SInter,C4H8O

r overa

ll,C

4H

8O

(1

06m

ol (g

cat. s

)1)


roverall,C4H8O

0

20

40

60

80

100

Carb

on

sele

ctivitie

s (

%)

(a) (b) (c)H-MFI H-FAU H4SiW12O40

96

Figure S3.4. Parity plot for the predicted and measured rates for intramolecular C=C bond

formation during butanal (C4H8O) reactions on H-MFI [ , space velocity 0.0033-0.013 mol

butanal (mol H+ s)

–1], H-FAU [▲, space velocity 0.0074-0.030 mol butanal (mol H

+ s)

−1], and

H4SiW12O40 [○, space velocity 0.045-0.18 mol butanal (mol H+ s)

−1] at 573 K.

Parity plots for the predicted and measured rate ratios, ,C H O,C H O 4 84 8

1-tetralin

( )j jr r ,

,C H O,C H O 4 84 8

1-tetralin-ad

( )j jr r , and ,C H O,C H O 4 84 8

1-chd

( )j jr r , on H-FAU (573 K), are shown in Figures

S3.5a, S3.5b, and S3.5c, respectively. Here, ,C H O4 8jr , ,C H O4 8 -tetralinjr ,

,C H O4 8 -tetralin-adjr , and

,C H O4 8 -chdjr are the rates for reaction j (j=Inter, Intra, Dehy, or Tish) in the reactant feed of C4H8O,

C4H8O-tetralin, C4H8O-tetralin-adamantane, and C4H8O-cyclohexadiene, respectively. The

predicted values for the rate ratios were determined via Equations S3.15a-S3.15c:

,C H O

,

,C H O

4 8

4 8

-tetralin

tetralin tetralin1

j

jj

rP

r (S3.15a)

,C H O

,

,C H O

4 8

4 8

-tetralin-ad

tetralin-ad tetralin1

j

jj

rP

r (S3.15b)

,C H O

,

,C H O

4 8

4 8

-chd

chd chd1

j

jj

rP

r (S3.15c)

0 20 40 60 800

20

40

60

80

r Intr

a,C

4H

8O

pre

dic

ted

(1

0-5

mo

l (m

ol H

+ s

)-1)

rIntra,C4H8O

measured (10-5

mol (mol H+ s)

-1)

0.00 0.05 0.10 0.15 0.200

1

2

3

4










8O

-C10H

12/r

C4H

8O


97

which are also Equations 3.5a-3.5c in Section 3.3.2. Here the proportionality constants, ,tetralinj ,

,tetralin-adj , and ,chdj , are listed in Table 3.2 of Section 3.3.2. Similarly, the parity plots for the

predicted and measured rate ratios ,C H O,C H O 4 84 8

1-

( )j jdr r

on H-MFI and H4SiW12O40 are shown in

Figures S3.6 and S3.7, respectively. Here d represents tetralin or cyclohexadiene (chd).


1-tetralin

( )j jr r ,

(b) ,C H O,C H O 4 84 8

1-tetralin-ad

( )j jr r , or (c) ,C H O,C H O 4 84 8

1-chd

( )j jr r ] for rates of butanal reactions in (a)

C4H8O-tetralin (,C H O4 8 -tetralinjr ), (b) C4H8O-tetralin-adamantane (

,C H O4 8 -tetralin-adjr ), or (c) C4H8O-

cyclohexadiene (,C H O4 8 -chdjr ) feed mixtures to those in C4H8O feed ( ,C H O4 8jr ) for intermolecular

C=C bond formation (Pathway 1, ), intramolecular C=C bond formation (Pathway 2, ○),

isomerization-dehydration (Pathway 3, ∆), and Tishchenko esterification-ketonization (Pathway

4, ■) on H-FAU at 573 K [subscript j=Inter, Intra, Dehy, or Tish, which denote inter- or


ketonization, respectively; space velocity 0.0074 mol butanal (mol H+ s)

-1, adamantane (if

added)=4-8 Pa].

0 1 2 3 40

1

2

3

4

r j,C

4H

8O

-ch

d (

r j,C

4H

8O

)1 p

redic

ted

rj,C4H8O-chd (rj,C4H8O)1

measured

0 1 2 3 40

1

2

3

4

r j,C

4H

8O

-te

tralin

-ad (

r j,C

4H

8O

)1 p

red

icte

d

rj,C4H8O-tetralin-ad (rj,C4H8O)1

measured

0 1 2 30

1

2

3

r j,C

4H

8O

-te

tralin

(r j,

C4

H8

O)

1 p

red

icte

d

rj,C4H8O-tetralin (rj,C4H8O)1

measured

(b)(a) (c)

0.00 0.05 0.10 0.15 0.200

1

2

3

4










8O

-C10H

12/r

C4H

8O


98


1-tetralin

( )j jr r

or (b) ,C H O,C H O 4 84 8

1-chd





(Pathway 2, ○), isomerization-dehydration (Pathway 3, ∆), and Tishchenko esterification-

ketonization (Pathway 4, ■) on H-MFI at 573 K [subscript j=Inter, Intra, Dehy, or Tish, which

denote inter- or intramolecular C=C bond formation, isomerization-dehydration, or Tishchenko

esterification-ketonization, respectively; space velocity 0.0033 mol butanal (mol H+ s)

-1].

0.0 0.5 1.0 1.5 2.00.0

0.5

1.0

1.5

2.0

r j,C

4H

8O

-ch

d (

r j,C

4H

8O

)1 p

redic

ted


measured

0.0 0.5 1.0 1.50.0

0.5

1.0

1.5

r j,C

4H

8O

-te

tra

lin (

r j,C

4H

8O

)1 p

redic

ted


measured

(b)(a)

0.00 0.05 0.10 0.15 0.200

1

2

3

4










8O

-C10H

12/r

C4H

8O


99


1-tetralin

( )j jr r

or (b) ,C H O,C H O 4 84 8

1-chd





(Pathway 2, ○), isomerization-dehydration (Pathway 3, ∆), and Tishchenko esterification-

ketonization (Pathway 4, ■) on H4SiW12O40 at 573 K [subscript j=Inter, Intra, Dehy, or Tish,

which denote inter- or intramolecular C=C bond formation, isomerization-dehydration, or

Tishchenko esterification-ketonization, respectively; space velocity 0.045 mol butanal (mol H+

s)-1

].

Parity plots for the predicted and measured rates for tetralin-to-alkanal transfer hydrogenation

(,C H O2TH -tetralinn n

r , Eqn. 3.9, n=3-6) during CnH2nO-tetralin reactions on H-FAU are shown in

Figure S3.8. Here the predicted ,C H O2TH -tetralinn n

r values were determined via Equation S3.16:

,C H O ,C H O2 2TH TH-tetralin -tetralin tetralinn nn nr k P (S3.16)

which is also Equation 3.9 in Section 3.3.3. Here the rate constants for tetralin-to-alkanal transfer

hydrogenation, ,C H O2TH -tetralinn n

k , are summarized in Figure 3.4 of Section 3.3.3.

0.0 0.5 1.0 1.50.0

0.5

1.0

1.5

r j,C

4H

8O

-te

tra

lin (

r j,C

4H

8O

)1 p

red

icte

d


measured

0.0 0.5 1.0 1.5 2.0 2.50.0

0.5

1.0

1.5

2.0

2.5

r j,C

4H

8O

-ch

d (

r j,C

4H

8O

)1 p

red

icte

d


measured

(b)(a)

0.00 0.05 0.10 0.15 0.200

1

2

3

4










8O

-C10H

12/r

C4H

8O


100

Figure S3.8. Parity plots for the predicted and measured rates for tetralin-to-alkanal transfer

hydrogenation (,C H O2TH -tetralinn n

r , Eqn. 3.9, n=3-6) during CnH2nO-tetralin reactions (▲, □, ,

and for n=3, 4, 5, and 6, respectively) on H-FAU [573 K, space velocity 0.0074 mol CnH2nO

(mol H+ s)

−1, 1.1 kPa CnH2nO, 0.08-0.16 kPa tetralin].

Parity plots for the predicted and measured rates for cyclohexadiene-to-butanal transfer

hydrogenation (,C H O4 8TH -chd

r , Eqn. 3.9) during on H-FAU, H-MFI, and H4SiW12O40 at 573 K are

shown in Figure S3.9. Here the predicted ,C H O4 8TH -chd

r values were determined via Equation

S3.17:

,C H O ,C H O2 2TH TH-chd -chd chdn nn nr k P (S3.17)

which is also Equation 3.9 in Section 3.3.3. Here the rate constants for cyclohexadiene-to-

butanal transfer hydrogenation on different catalysts, ,C H O4 8TH -chd

k , are summarized in Table 3.2

of Section 3.3.2.

0.0 0.2 0.4 0.6 0.8 1.00.0

0.2

0.4

0.6

0.8

1.0

r TH

,CnH

2nO

-tetr

alin

pre

dic

ted (

10

-3 m

ol (m

ol H

+ s

)-1)

rTH,CnH2nO-tetralin

measured (10-3

mol (molH+ s)-1

)

0.0 0.2 0.4 0.6 0.8 1.0 1.20.0

0.2

0.4

0.6

0.8

1.0

r TH

,Cn

H2

n-t

etr

alin

(10

-3 m

ol (m

ol H

+ s

)-1)

PTetralin (kPa)

0.0 0.2 0.4 0.6 0.8 1.0 1.20.0

0.2

0.4

0.6

0.8

1.0

r TH

,Cn

H2

n-t

etr

alin(1

0-3

mo

l (m

ol H

+ s

)-1)

PTetralin (kPa)

101

Figure S3.9. Parity plots for the predicted and measured rates for cyclohexadiene-to-butanal

transfer hydrogenation (,C H O4 8TH -chd

r , Eqn. 3.9) on H-FAU (▲), H-MFI (■), and H4SiW12O40 (●)

at 573 K [space velocity=0.0074, 0.0033, and 0.045 mol butanal (mol H+ s)

−1 for H-FAU, H-

MFI, and H4SiW12O40, respectively, 1.1 kPa butanal, 0.03-0.15 kPa cyclohexadiene].

3.6.6. Characterizations of the Brønsted and Lewis acid sites

The Brønsted and Lewis acid sites on H-MFI, and H-FAU, and H4SiW12O40 catalysts were

characterized by pyridine titration and by an infrared spectroscopic study of pyridine adsorption.

Table S3.2 summarizes the amounts of Brønsted and Lewis acid sites on these catalysts. The

total amount of acid sites was determined based on the pyridine uptake during the pyridine

titration at 473 K. The percentage of Brønsted and Lewis acid sites were determined based on the

infrared absorption bands at 1545 and 1455 cm−1

, respectively, measured during pyridine

adsorption on the catalysts at 473 K [49].

0.0 0.2 0.4 0.60.0

0.2

0.4

0.6

r TH

,C4

H8

O-c

hd p

red

icte

d (

10

-3 m

ol (m

ol H

+ s

)-1)

rTH,C4H8O-chd

measured (10-3

mol (molH+ s)-1

)

102

Table S3.2. The amounts of Brønsted and Lewis acid sites on H-MFI, and H-FAU, and

H4SiW12O40 catalysts

H-MFI H-FAU H4SiW12O40

Total acid sites (μmol gcat.−1

)a 1120 506 169

Brønsted acid percentageb 84 % 81 % 71 %

Lewis acid percentageb 16 % 19 % 29 %

Brønsted acid sites(μmol gcat.−1

) 944 410 120

Lewis acid sites (μmol gcat.−1

) 176 96 49

aThe amounts of total acid sites were determined by pyridine titration at 473 K; bThe percentage of Brønsted and Lewis acid sites were determined based on the infrared spectra of pyridine adsorption at 473 K

[49].

103

Chapter 4 Kinetic Requirements of Solid Brønsted Acid Catalyzed

Transfer Hydrogenations of Aldehyde

Abstract

Solid Brønsted acids in confined environment catalyze transfer hydrogenation of aldehydes

(CnH2nO, n=3-6) via a kinetically-relevant hydride transfer step between hydride donors (e.g.,

cyclohexadiene, tetralin, cyclohexene, or 3-methyl-1-pentene) and protonated aldehydes. The

hydride transfer occurs preferentially on partially confined H+ sites and involves a bimolecular

transition state consisted of the carbenium ions of the hydride donor and acceptor, both sharing

the hydride ion. The hydride ion affinity difference between these two donor and acceptor

fragments relates directly to the heat of hydride transfer reaction; thus it is a kinetic descriptor

that dictates the overall transfer hydrogenation rates through the classical Brønsted-Evans-

Polanyi relation, which correlates the heat of hydride transfer reaction to its barrier.

4.1. Introduction

Transfer hydrogenation is ubiquitous across homogeneous and heterogeneous catalysis. The

reaction typically occurs on transition metals or metal complexes, but Brønsted acids may also

catalyze the asymmetric transfer hydrogenation of imines, aldehydes, or ketones [1-4]. In

heterogeneous systems, solid Brønsted acids (H+; H-MOR zeolite [5]) protonate alkenes that

sequentially accept hydride ions from alkanes [5-7]. Here, I show that solid Brønsted acids

confined in pores and cages of molecular dimensions may also protonate aldehydes and initiate

hydride transfer effectively. The hydride transfer is kinetically-relevant in aldehyde transfer

hydrogenation and is mechanistically analogous to those found during heterogeneous,

homogeneous, and enzymatic reductions of: (1) quinoline protonated by chiral phosphoric acid

accepts a hydride ion from substituted dihydropyridine (Scheme 4.1a) [4], (2) pyruvate

protonated by lactate dehydrogenase (LDH) accepts a hydride ion from nicotinamide adenine

dinucleotide (NADH) (Scheme 4.1b) [8], and (3) alkene protonated by acidic zeolite, accepts a

hydride ion from an alkane (Scheme 4.1c) [7]. These reactions all begin with an initial

104

protonation (e.g., C=O to C-O-H+, C=NH to C-NH2

+, and C=C to C-CH

+ ions) followed by

hydride transfer and share similar transition state structures (Scheme 4.1). I further establish the

mechanistic details and pinpoint the catalytic requirements for aldehyde transfer hydrogenation,

in terms of the thermochemical properties of reactants and local H+ site confinements. I show

that (1) the difference in the hydride ion affinities (∆𝐻𝐼𝐴RDH+−C𝑛H2𝑛OH+) between the carbenium

ions of the hydrogen donor (RDH+) and protonated hydrogen acceptor (CnH2nOH

+) is a kinetic

descriptor that dictates the hydride transfer rates and (2) local H+ site confinements solvate the

bimolecular transition state formed between the H-donor and protonated aldehyde (Scheme 4.1d)

and thus promote the hydride transfer events.

Scheme 4.1. Pathways of transfer hydrogenation of (a) protonated quinoline catalyzed by chiral

Brønsted acid (chiral phosphoric acid) [4], (b) pyruvate catalyzed by lactate dehydrogenase [8],

(c) protonated alkene catalyzed by solid Brønsted acid sites (e.g., H-MOR zeolite [5], and H3Si-

OH-AlH2-O-SiH3 cluster model [7]), and (d) protonated aldehyde catalyzed by solid Brønsted

acid sites (R, R1, and R2 denote alkyl groups). In each of these cases, the reaction involves

105

protonation of the reactant followed by hydride transfer (from a hydride donor) (H-acceptor and

H-donor denote hydride acceptor and hydride donor, respectively).

4.2. Experimental


H-FAU zeolite samples were prepared by treating their H+ form (780 m

2 g

-1, Si/Al atomic

ratio=15, CBV720, Zeolyst), in flowing dry air (0.6 cm3 g

-1 s

-1, zero grade, Linde), heated to 873

K at 0.0167 K s-1

and holding for 4 h.

4.2.2 Rate and selectivity assessments

Catalytic rates of aldehyde (e.g. propanal, butanal, pentanal, and hexanal) reactions were

measured under differential conditions in a fixed bed microcatalytic quartz reactor (9.5 mm inner

diameter) with plug-flow hydrodynamics. The reactor was contained within a resistively heated

furnace controlled by a digital feedback temperature controller. 100 mg of zeolites were

supported on a quartz frit and the catalyst bed temperature was recorded using a K-type

thermocouple placed at the center (in both the axial and radial directions) of the catalyst bed.

Catalysts were treated in-situ under flowing He (Grade 5.0, Linde, 8.3 cm3 gcat.

-1 s

-1) at 0.167 K s

-

1 to the reaction temperature (573 K) prior to catalytic rate measurements. Aldehyde reactants

[butanal (Sigma Aldrich, purity grade, ≥99%, CAS# 123-72-8), propanal (Sigma Aldrich, 97%,

CAS# 123-38-6), pentanal (Sigma Aldrich, 97%, CAS# 110-62-3), or hexanal (Sigma Aldrich,

98%, CAS# 66-25-1)] and hydride donor reactants [tetralin (Sigma Aldrich, 99%, CAS# 119-64-

2), cyclohexadiene (Sigma Aldrich, 97 %, CAS# 592-57-4), cyclohexene (Sigma Aldrich, 99 %,

CAS#110-83-8 ), 3-methyl-1-pentene (Sigma Aldrich, 99%, CAS# 760-20-3), or cyclohexane

(Sigma Aldrich, 99.5 %, CAS#110-82-7 )] were introduced via gas-tight syringes (either 5 cm3

Hamilton Model 1005 or 1 cm3 SGE Model 008025 for aldehydes, and 0.25 cm

3 SGE Model

006230 for hydride donors), which were mounted on syringe infusion pump (KD Scientific,

LEGATO 100), into two vaporization zones heated to the boiling points of the respective

reactants at atmospheric pressure, within which liquid reactants were evaporated and mixed with

a He (Linde, Grade 5.0, 8.3 cm3 gcat.

−1 s

−1) purge stream. The mixture was fed to the reactor via

heated transfer lines held isothermally at 473 K. Chemical species in the reactor effluent stream

106

were quantified using an on-line gas chromatograph (Model 7890A, Agilent) and mass

spectrometer (Model 5975C, Agilent), GC-MS, equipped with a 10-port sampling valve

containing two sample loops of 250 μl each. The samples contained in the gas sampling loops

were analyzed by chromatographic separation using two capillary columns (HP-5, Agilent,

19091J-413, 30 m, 0.32 mm ID and HP-5MS, Agilent, 190091S-433, 30 m, 0.25 mm ID). The

HP-5 column is connected to thermal conductivity (TCD) and flame ionization (FID) detectors in

series and the HP-5MS column to the MS detector.

4.2.3. Infrared spectroscopic study

Infra-red (IR) spectroscopic studies were carried out with a customized, stainless steel in-situ

transmission IR cell equipped with CaF2 windows and capable of operating between 298-773 K.

The cell was mounted in a Bruker Vertex 70 spectrometer equipped with a mercury cadmium

telluride (MCT) detector. Powder samples were pressed into a self-supporting wafer, ca. 10 mm

in diameter and less than 0.5 mm in thickness, mounted on a sample holder inside the cell.

Infrared spectra were acquired in the transmission mode at a resolution of 2 cm-1

and 16 scans

per spectrum. Prior to the adsorption studies, samples were heated at 0.167 K s-1

to 573 K and

kept at 573 K in flowing helium (0.83 cm3

s-1

) for 30 min. Adsorptions of butanal, pyridine,

butanol, and hydrogen donors (cyclohexadiene, tetralin, cyclohexene, 3-methyl-1-pentene, and

cyclohexane) were carried out at 373 K. The chemicals were introduced by infusing them via a

gas tight syringe (0.25 cm3 SGE Model 006230) mounted on syringe infusion pump into a

vaporization zone, which was maintained at their respective boiling points. In the vaporization

zone, the liquid reactants were mixed with a He (0.83 cm3 s

−1) purge stream and fed to the IR

cell via heated transfer lines held isothermally at 373 K, which were connected to an on-line gas

chromatograph to quantify the concentration of reactants.

107


4.3.1. Kinetic and infrared spectroscopic studies on aldehyde transfer hydrogenation by hydrocarbons on Brønsted acid sites

The protonated aldehydes (CnH2nO, n=3-6) undergo aldol condensation reactions, forming larger,

unsaturated aldehydes (C2nH4n-2O), which upon secondary aldol condensation, ring closure, and

dehydration [9], evolve alkyl cycloalkadienes or aromatics (Scheme 4.2). In parallel, the

protonated aldehydes (CnH2nO) may also undergo transfer hydrogenation followed by

dehydration, forming alkenes while preserving their carbon backbone via:

Protonation: CnH2nO+H+CnH2nOH

+ (4.1a)

Hydride transfer: CnH2nOH+ +RDH2 CnH2n+1OH +RDH

+ (4.1b)

Dehydration: CnH2n+1OHCnH2n+H2O (4.1c)

H+

site regeneration: RDH+RD+H

+ (4.1d)

where RDH2 denotes a hydride donor (H-donor, e.g., RDH2=alkyl tetralin, alkyl cyclohexadiene).

Clearly, the transfer hydrogenation reaction requires hydrogen atoms from the H-donors.

Without an external hydrogen source, these donors must originate from the aromatic pools

produced from the parallel aldol condensation and secondary ring closure reactions shown in

Scheme 4.2.

108

Scheme 4.2. Pathways of aldehyde reactions that generate hydrogen donors (aromatics or their

precursors, labeled H-donor), which include aldol condensation and ring closure steps (R and R1-

R4 represent either an H or alkyl groups). The parallel pathways of aromatic transalkylation and

their products are not shown in the scheme for simplification purposes (the complete reaction

network is reported elsewhere[3]).

Alkene site-time-yields are linearly proportional to the pressure of aromatic fractions during

aldehyde (CnH2nO, n=3-6) reactions on H-MFI [3, 9], H-FAU [3], and H4SiW12O40[10] catalysts,

thus aromatics are the H-donors. The involvement of diverse aromatics, which range from C7 to

C16+ (see Fig. S4.2 in Sec. 4.6.2 for detailed distributions of the aromatic fraction during butanal

reactions on H-FAU, 573 K), as the hydrogen donors has complicated the rate assessments,

simply because the transfer hydrogenation occurs concomitantly with steps that produce the H-

donors within the complex reaction network. I eliminated this complication, by probing

specifically the transfer hydrogenation rates (Eqn. 4.1) with a series of H-donors

(RDH2=cyclohexadiene, tetralin, cyclohexene, 3-methyl-1-pentene, and cyclohexane) and

CnH2nO (n=3-5) aldehydes on H-FAU zeolites (Si/Al=15). I determined the transfer

hydrogenation rates between these H-donor-acceptor pairs from the overall alkene (CnH2n) site-

time-yields, after isolating the rate contributions from those enabled by the secondary products

(methods described in Sec. 4.6.3). As shown in Figure 4.1a, transfer hydrogenation

109

rates, 𝑟TH,C𝑛H2𝑛O−RDH2 (TH denotes transfer hydrogenation and subscript CnH2nO-RDH2 the feed

mixture), between butanal and a series of H-donors, increase proportionally with H-donor

pressure (𝑃RDH2) at 573 K:

𝑟TH,C𝑛H2𝑛O−RDH2=

𝑘TH,C𝑛H2𝑛O−RDH2𝑃RDH2𝐾 C𝑛H2𝑛O𝑃C𝑛H2𝑛O

1 + 𝐾 C𝑛H2𝑛O𝑃C𝑛H2𝑛O + 𝛼

(H+ sites) (Adsorbed aldehydes) (Other)

(4.2)

where 𝑘TH,C𝑛H2𝑛O−RDH2is the rate constant of hydrogen transfer, 𝑃C𝑛H2𝑛O is the aldehyde

pressure, and 𝐾 C𝑛H2𝑛O is the equilibrium constant of aldehyde adsorption. The term

𝐾 C𝑛H2𝑛O𝑃C𝑛H2𝑛O represents the surface concentration of the adsorbed CnH2nO and 𝛼 the

concentration of all other surface species (e.g., products and H-donors). Equation 4.2 becomes:

𝑟TH,C𝑛H2𝑛O−RDH2= 𝑘TH,C𝑛H2𝑛O−RDH2

𝑃RDH2 (4.3)

when adsorbed CnH2nO is the most abundant surface intermediate (𝐾 C𝑛H2𝑛O𝑃C𝑛H2𝑛O>>1+ 𝛼).

The slopes in Figure 4.1a give the rate constant values for 𝑘TH,C𝑛H2𝑛O−RDH2. This linear

dependence remains true for the transfer hydrogenation rates across all CnH2nO (n=3-6)

aldehydes, as they vary linearly with the pressure of H-donors produced from the secondary

reactions on H-MFI zeolites [3] and H4SiW12O40 clusters [3, 10].

110

Figure 4.1. (a) Rates (D4 8 2TH,C H O R H-r ) for butanal (C4H8O) transfer hydrogenation as a function of

H-donor pressure (D 2R HP ); (b)-(d) rate constants (

D4 8 2TH,C H O R H-k , Eqn. 4.3) for (b) butanal, (c)

propanal, and (d) pentanal transfer hydrogenation by various H-donors (RDH2), plotted as a

function of the hydride ion affinity difference ( + +D 2R H C H OH- n n

HIA , Eqn. 4.4) between the

carbenium ions of H-donor (RDH+, e.g., RDH

+=C10H11

+ for tetralin as the hydrogen donor) and

the protonated aldehydes (CnH2nOH+). The identities of H-donor are shown in the figure (573 K,

RDH2=cyclohexadiene, tetralin, cyclohexene, 3-methyl-1-pentene, or cyclohexane, H-FAU

(Si/Al=15)). The dash lines in (b)-(d) reflect the predicted reactivity trend of C6 H-donors

(cyclohexadiene, cyclohexene, and 3-methyl-1-pentene, which have similar molecular sizes)

based on + +D 2R H C H OH- n n

HIA .

111

I propose a mechanism in Scheme 4.3a, provide the kinetic evidence, and then describe the

catalytic requirements for transfer hydrogenation (Step 1), which occurs via kinetically-relevant

hydride ion (H−) transfer (see Sec. 4.6.4 for the discussion about the kinetically-relevant step).

The reaction begins with aldehyde protonation (as CnH2nOH+) and saturation of the Brønsted

acid sites as the most abundant surface intermediates, as confirmed from the complete

disappearance of the OH stretching vibration (Fig. S4.1a in Sec. 4.6.1). Next, an hydride ion

transfers from a hydrogen donor (RDH2) to a protonated aldehyde (CnH2nOH+) via a

carbocationic transition state [CnH2nOH+∙∙∙H

−∙∙∙RDH

+]

‡ (Step 1.1). At the transition state, the

carbenium ion of the H-donor (RDH+) and the protonated aldehyde (CnH2nOH

+) share a hydride

ion (H−). Decomposition of the [CnH2nOH

+∙∙∙H

−∙∙∙RDH

+]‡ transition state (Step 1.2) completes the

hydride ion transfer, leading to an alcohol (CnH2n+1OH) and a carbenium ion (RDH+; e.g.,

RDH+=C10H11

+, when tetralin is the H-donor). This hydride transfer mechanism and its transition

state structure resemble those between alkanes and alkenes[6, 7, 11] and between alkanes and

alkoxides[5] on acidic zeolites, proposed based on density functional theory (DFT) calculations

(Scheme 4.1c). The mechanism also mimics those in homogeneous or enzymatic reactions

(Schemes 4.1a and 4.1b), as described earlier.

After the kinetically-relevant step (Step 1, Scheme 4.3a), the carbenium ion RDH+ donates a H

+

to the zeolitic framework, forming the corresponding dehydrogenated product (RD, RD=C10H10

for tetralin) and regenerating the Brønsted acid site (Step 2). The alcohol (CnH2n+1OH) undergoes

sequential rapid, kinetically-irrelevant acid-catalyzed dehydration that evolves alkene (CnH2n)

(Step 3), as confirmed from the much faster butene formation rates for 1-butanol dehydration

(Step 3) than butanal transfer hydrogenation (Steps 1-3, by >38 and >22 times on H-MFI [3] and

H-FAU, respectively) under identical conditions (1.1 kPa, 573 K).

Scheme 4.3b shows a Born-Haber thermochemical cycle construct used for estimating the heat of

hydride transfer reaction. The heat of reaction (∆r𝐻TH) decomposes into the hydride ion affinity

for the carbenium ion of the hydride donor (RDH+), 𝐻𝐼𝐴RDH+, and for the protonated aldehyde

(CnH2nOH+), 𝐻𝐼𝐴C𝑛H2𝑛OH+ . Thus, it equals the hydride ion affinity difference between the H-

donor and acceptor pair (∆𝐻𝐼𝐴RDH+−C𝑛H2𝑛OH+):

112

∆r𝐻TH = 𝐻𝐼𝐴RDH+ − 𝐻𝐼𝐴C𝑛H2𝑛OH+ = ∆𝐻𝐼𝐴RDH+−C𝑛H2𝑛OH+ (4.4)

The higher exothermicity and thus more negative hydride ion affinity difference,

∆𝐻𝐼𝐴RDH+−C𝑛H2𝑛OH+ , lead to smaller activation enthalpy ( ∆𝐻‡ ) for the formation of

[CnH2nOH+∙∙∙H

−∙∙∙RDH

+]‡ transition state (Step 1.1, Scheme 4.3a), according to Brønsted-Evans-

Polanyi relation.

Scheme 4.3. (a) Catalytic steps (Steps 1-3) for solid Brønsted acid catalyzed transfer

hydrogenation of aldehydes (CnH2nO) by hydrogen donors (RDH2) and the proposed hydride

transfer transition state; (b) the heat of reaction for the kinetically-relevant hydride transfer step

( THrH , Step 1), interpreted using a Born-Haber thermochemical construct ( +DR H

HIA and

+2C H OHn n

HIA denote the hydride ion affinities of the H-donor carbenium ion and protonated

aldehyde, respectively).

113

The direct relation between the hydride ion affinity differences and rate constants for transfer

hydrogenation was first confirmed by varying the chemical identity and hydride ion affinity of

the H-donors (𝐻𝐼𝐴RDH+). Butanal transfer hydrogenation rate constants ( 𝑘TH,C4H8O−RDH2) for a

variety of H-donors are plotted in Figure 4.1b as a function of the hydride ion affinity

difference, ∆𝐻𝐼𝐴RDH+−C4H8OH+ . Only the H-donor and acceptor pairs with negative

∆𝐻𝐼𝐴RDH+−C4H8OH+ values (RDH2= cyclohexadiene, tetralin, or cyclohexene) were found to be

reactive. Among the reactive H-donors, higher rate constants (𝑘TH,C4H8O−RDH2) were found for

the donor-acceptor pairs with larger and more negative ∆𝐻𝐼𝐴RDH+−C4H8OH+ values. In contrast,

no reactivity was detected for an H-donor and acceptor pair with positive or near zero

∆𝐻𝐼𝐴RDH+−C4H8OH+values (RDH2=cyclohexane or 3-methyl-1-pentene). Similar reactivity trends

were also found for propanal and pentanal transfer hydrogenation (Figs. 4.1c and 4.1d). Previous

DFT studies[5] show that the hydride transfer between an alkane H-donor (isobutane, propane,

and ethane) and ethyl alkoxide H-acceptor in H-MOR forms similar transition state structures

and the corresponding activation barrier follows the order of C2H6>C3H8>i-C4H10, because larger

or more substituted carbenium ions form more stable transition states [5]. This reactivity order is

consistent with the trend derived from the hydride ion affinities of their respective carbenium

ions (C2H5+, C3H7

+ and i-C4H9

+): 𝐻𝐼𝐴C2H5

+ (1131 kJ mol−1

)> 𝐻𝐼𝐴C3H7+ (1049 kJ mol

−1)>

𝐻𝐼𝐴𝑖−C4H9+ (990 kJ mol

−1) (see Sec. 4.6.8 for the determination of 𝐻𝐼𝐴RDH+ values). A higher

hydride ion affinity of the H-donor decreases their thermodynamic tendency to donate a hydride

ion. Previous studies reported these trends, but did not draw a direct connection between the

hydride ion affinity differences (Eqn. 4.4) and the transfer hydrogenation rate constants.

The direct correlation between transfer hydrogenation rate constants (𝑘TH,C𝑛H2𝑛O−RDH2) and

hydride ion affinity differences (∆𝐻𝐼𝐴RDH+−C4H8OH+ , Eqn. 4.4) is shown conclusively with

infrared spectroscopic studies. First, 10 Pa butanal was introduced to H-FAU at 373 K to attain a

butanal fractional coverage ( 𝜃C4H8O,0 ) between 0.8 and 0.12 (Table S4.1 in Sec. 4.6.5),

confirmed from the disappearance of the hydroxyl stretching bands at 3625 cm-1

and 3560 cm-1

(time-resolved spectra upon butanal adsorption and desorption are shown in Fig. S4.3a in Sec.

4.6.5) and the concomitant evolvement of the C=O band of adsorbed butanal at 1675 cm-1

(Figs.

4.2a-4.2b and Fig. S4.5 in Sec. 4.6.6). Subsequently, butanal feed was removed and replaced

with helium flow, during which a portion of the protonated butanals was desorbed from the H+

114

sites, as confirmed from the decrease in the C=O band intensity (Fig. 4.2a). Most H+ sites

(~80 %, Fig. S4.3b in Sec. 4.6.5) were recovered in that process, whereas the remaining sites

(~20 %) were occupied by the heavier, condensation products (e.g., 2-ethyl-2-hexenal and

aromatics).

In separate cases, a H-donor (10-15 Pa of cyclohexadiene, tetralin, cyclohexene, or cyclohexane)

was introduced to H-FAU zeolites following butanal adsorption. In the presence of H-donors, the

C=O absorption band at 1675 cm-1

decreased much more rapidly (see Fig. 4.2b and Fig. S4.5 in

Sec. 4.6.6 for time-resolved infrared spectra) than those without a H-donor (Fig. 4.2a), because a

portion of the protonated butanals undergoes transfer hydrogenation. H-donors did not displace

the protonated butanals and cause the decay in C=O band intensity, despite the fact that some of

these H-donors exhibit higher proton affinities (PA) than butanal (e.g., 837 kJ mol−1

for

cyclohexadiene and 810 kJ mol−1

for tetralin vs. 793 kJ mol−1

for butanal).[12] In fact,

incorporating these H-donors (8-16 Pa) during steady-state butanal (1.1 kPa) reactions did not

perturb the parallel intermolecular C=C bond formation reaction (Table S4.2 in Sec. 4.6.7), an

indication that H-donors did not affect the butanal surface coverages.

Figure 4.2c shows the decay of C=O bands or the equivalent butanal coverages (𝜃C4H8O, see Sec.

4.6.6 for methods of determining the coverages) as a function of time when exposing the H-FAU

sample to different H-donors; the rates of 𝜃C4H8O decay in the presence of the various H-donors

(RDH2) follow the order of: cyclohexadiene ( ) > tetralin ( ) > cyclohexene ( ) > cyclohexane

() ≈ without H-donor (▲). This decay in 𝜃C4H8O is a combined result of butanal desorption and

butanal consumption via transfer hydrogenation, captured by:

𝑑𝜃C4H8O

𝑑𝑡= −(𝑘des−IR𝜃C4H8O + 𝑘TH−IR,RDH2

𝑃RDH2𝜃C4H8O ) (4.5)

where t represents the duration of H-donor addition, 𝑘des−IR is the rate constant for butanal

desorption, 𝑘TH−IR,RDH2 is the rate constant for transfer hydrogenation, and 𝑃RDH2

is the partial

pressure of H-donors. Integrating Equation 4.5, together with the initial butanal coverage

(𝜃C4H8O,0), gives an expression that captures the time-dependent butanal coverages:

𝜃C4H8O = 𝜃C4H8O,0exp [−(𝑘des−IR + 𝑘TH−IR,RDH2𝑃RDH2

)𝑡] (4.6)

115

The rate constant for butanal desorption, 𝑘des−IR, was determined from non-linear regression of

the decay in butanal coverages during butanal desorption (▲, without H-donor, 𝑃RDH2=0, Figure

4.2c) against Equation 4.6. Similarly, the rate constants for butanal transfer hydrogenation using

the various H-donors, 𝑘TH−IR,RDH2, were determined from non-linear regression of the respective

data in Figure 4.2c against Equation 4.6 and with the 𝑘des−IR value. As shown in Figure 4.2d,

these rate constants determined from the decay in butanal coverages, 𝑘TH−IR,RDH2, are directly

proportional to those measured during steady-state reactions in a micro-catalytic fixed bed

reactor. Their linear relation, taken together with the direct correlation between rate constants

and hydride ion affinity differences in Figures 4.1b-4.1d, confirms the kinetic relevance of

hydride ion transfer and the involvement of H-donors.

116

Figure 4.2. (a) and (b) Time-resolved infrared spectra upon exposure of H-FAU (Si/Al=15) to 10

Pa butanal followed by purging (a) in He or (b) in 15 Pa tetralin at 373 K; (c) coverages of

butanal on the H+ sites of H-FAU (

4 8C H O ) as a function of time upon purging with He or

introducing various H-donors at 373 K (line: fitted profiles against Eqn. 4.6); (d) butanal transfer

hydrogenation rate constants derived from in-situ infrared absorption spectroscopy (D 2TH-IR,R Hk ,

373 K) plotted against the transfer hydrogenation rate constants measured with steady-state

micro-catalytic flow reactor (D4 8 2TH,C H O R H-k , 573 K) with various H-donors (H-donors and

+ +D 4 8R H C H OH-

HIA values are indicated in the figures)

117

4.3.2. Effects of carbon chain length on the aldehyde transfer hydrogenation

The direct connection between the hydride transfer rate constants (𝑘TH,C𝑛H2𝑛O−RDH2) and the

hydride ion affinity differences (∆𝐻𝐼𝐴RDH+−C𝑛H2𝑛OH+) was also confirmed separately by varying

the chemical identity and hydride ion affinity of the H-acceptor (protonated aldehyde, CnH2nOH+)

while using tetralin as the H-donor (𝐻𝐼𝐴C10H11+=934.1 kJ mol

−1). As the aldehyde chain length

increases from C3 to C6, the hydride ion affinity of the protonated aldehyde CnH2nOH+

(𝐻𝐼𝐴C𝑛H2𝑛OH+) decreases from 956.6 to 941.1 kJ mol−1

and the hydride ion affinity difference

(∆𝐻𝐼𝐴C10H11+−C𝑛H2𝑛OH+, Eqn. 4.4) becomes commensurately more negative. As shown in Figure

4.3, smaller aldehydes with more negative hydride ion affinity differences are more reactive in

transfer hydrogenation by aromatic H-donors on H4SiW12O40 polyoxometalate clusters

(𝑘TH, C𝑛H2𝑛O, ▲) [10]. Similar trends remain on H-FAU, on which the rate constants for CnH2nO

(n=4-6) transfer hydrogenation by tetralin ( 𝑘TH,C𝑛H2𝑛O−C10H12) increase as the hydride ion

affinity differences become more negative (Fig. 4.3, ) [3], except for propanal. Propanal is less

reactive than butanal in transfer hydrogenation, despite its more negative hydride ion affinity

difference (−22.5 vs. −13.7 kJ mol−1

for propanal and butanal, respectively) [3]. I provide next a

hypothesis on this deviation and then confirm our hypothesis by using a different H-donor.

Confined environment is known to solvate and stabilize the hydride transfer transition state

formed between alkane and protonated alkene, as confirmed previously from alkane cracking in

SAPO-41, H-MFI, and H-Y zeolites [13] and from dimethyl ether homologation in acidic

zeolites (FAU, BEA, and MFI), SiO2-Al2O3, and H4SiW12O40/SiO2 catalysts [14]. Similar effects

of confinement are found during butanal transfer hydrogenation on H-MFI, H-FAU, and

H4SiW12O40 catalysts [3]. The deviation of propanal from the expected reactivity trend in Figure

4.3 () is caused apparently by its smaller transition state [C10H11+∙∙∙H

−∙∙∙C3H6OH

+]‡ for hydride

transfer [structure shown in Fig. 4.3(ii)] that is fitted loosely within the FAU supercage (11.8 Å)

and thus is solvated to a lesser extent, because of the smaller kinetic diameter of propanal (4.7 Å)

than butanal (5.0 Å). Such solvation effects are also found during hydride transfer in n-pentane

cracking, as an hydride ion is being transferred from a n-C5H12 onto C2H5+, C3H7

+, or C4H9

+

carbenium ion [13], and during dimethyl ether homologation, as an hydride ion is being

transferred from an alkane onto a C7H15+ carbenium ion [14], In both of these cases, zeolites with

appropriate pore sizes solvate and stabilize the bimolecular transition states (e.g., CIT-1 and

118

MCM-68 zeolites for n-pentane cracking [13] and FAU and BEA zeolites for dimethyl ether

homologation[14]), thus they are more effective in catalyzing the hydride transfer than those

with either smaller or larger pore dimensions.

To further confirm the confinement effects, I varied the size of the H-donor. When using

cyclohexadiene (C6H8, kinetic diameter 5.9 Å[15]), a smaller H-donor, the rate constants initially

increased as the size of aldehyde decreased (from hexanal to pentanal), as expected because of

the more negative hydride ion affinity difference (∆𝐻𝐼𝐴C6H7+−C𝑛H2𝑛OH+ , Fig. 4.3, ). Both

propanal and butanal, however, exhibited lower rate constants 𝑘TH,C𝑛H2𝑛O−C6H8 than pentanal,

contradicting the reactivity trend predicted based solely on ∆𝐻𝐼𝐴C6H7+−C𝑛H2𝑛OH+ , because of

their smaller transition states [CnH2nOH+∙∙∙H

−∙∙∙C6H7

+]

‡ [structure shown in Fig. 4.3(i)] and thus

lesser extents of solvation by the FAU structures via van der Waals interactions. Cyclohexadiene

(C6H8, kinetic diameter 5.9 Å[15]) as a smaller H-donor than tetralin (C10H12, ~6.2 Å[16]) forms

smaller transition states with both propanal (4.7 Å) and butanal (5.0 Å)

([CnH2nOH+∙∙∙H

−∙∙∙C6H7

+]

‡, n=3 or 4); these transition states are unable to fill the FAU

supercages, thus their reactivities deviate from those predicted based solely on

∆𝐻𝐼𝐴C6H7+−C𝑛H2𝑛OH+. Among the H-donors examined in this work, tetralin (C10H12, ~6.2 Å) has

a larger molecular size than the C6 hydrocarbons [cyclohexadiene (C6H8, ~5.9 Å), cyclohexene

(C6H10, ~5.9 Å), and 3-methyl-1-pentene (C6H12, ~5.6 Å)], and this size difference leads to

different extents of transition state solvation. During the transfer hydrogenation of a smaller H-

acceptor (propanal or butanal, 4.7-5.0 Å), the [CnH2nOH+∙∙∙H

−∙∙∙RDH

+]‡ (n=3 or 4) transition

states formed with the smaller C6 H-donors (RDH2=cyclohexadiene, cyclohexene, or 3-methyl-1-

pentene) are commensurately smaller and fitted loosely in the FAU cage than those formed with

tetralin (RDH2=tetralin) as the larger donor. As a result, transfer hydrogenation with tetralin gives

a higher rate constant than the trend predicted from the C6 H-donors predicted solely based on

∆𝐻𝐼𝐴RDH+−C𝑛H2𝑛OH+ (dotted line in Figs. 4.1b and 4.1c). In the contrasting case of a larger H-

acceptor (pentanal, 5.7 Å, Fig. 4.1d), the [C5H10OH+∙∙∙H

−∙∙∙RDH

+]

‡ transition state formed with

the C6 H-donors can fill completely the FAU cage and thus are stabilized by solvation. The

transition state formed between pentanal and tetralin is “oversized” and destabilized by steric

constraints, making tetralin less reactive than the predicted trend for the series of C6 H-donors

(dotted line in Fig. 4.1d). These results in Figures 4.3 and 4.1 suggest that transfer hydrogenation

119

rates are dictated by both the hydride ion affinity difference, a reactant property, and the size of

the transition state, the latter through their solvation effects.

Figure 4.3. The rate constants for aldehyde (CnH2nO, n=3-6) transfer hydrogenation with

cyclohexadiene (C6H8, 2 6 8TH,C H O C H-n n

k , , transition state depicted in Fig. 4.3(i)) or tetralin

(C10H12, 2 10 12TH,C H O C H-n n

k , , from [3], transition state depicted in Fig. 4.3(ii)) as the H-donor on

H-FAU zeolite (573 K), plotted as a function of hydride ion affinity difference

( + +10 11 2C H C H OH- n n

HIA or + +76 2C H C H OH- n n

HIA ); the rate constant for aldehyde transfer hydrogenation

with aromatic products as the H-donors (2,C H OTH n n

k , ▲, from[10], transition state depicted in Fig.

4.3(iii)) on H4SiW12O40 (573 K), plotted as a function of the hydride ion affinity difference

( + +10 11 2C H C H OH- n n

HIA ).

120

4.4. Conclusion

Protonated aldehydes (CnH2nOH+, n=3-6) on solid Brønsted acid sites undergo transfer

hydrogenation with hydrocarbons as the H-donors, leading to alkene (CnH2n) formation. This

mechanism is confirmed by incorporating H-donor into the aldehyde reaction on H-FAU zeolite,

which promotes the consumption of protonated carbonyl group and increases the rate of alkene

formation. This acid catalyzed transfer hydrogenation proceeds via hydride ion transfer that

involves a bimolecular transition state formed from the H-donor and acceptor pair sharing a

hydride ion. The rates are dictated by the hydride ion affinity difference between the carbenium

ions of the hydride donor and acceptor pair. The transition state is stabilized through solvation by

the local environment of zeolites: smaller transition state that fits loosely in the cage is stabilized

to a lesser extent, thus leading to a lower reactivity for hydride transfer.

4.5. References

[1] S. Rossi, M. Benaglia, E. Massolo, L. Raimondi, Organocatalytic strategies for

enantioselective metal-free reductions, Catal. Sci. Tech., 4 (2014) 2708-2723.

[2] T. Prasomsri, R.E. Galiasso Tailleur, W.E. Alvarez, T. Sooknoi, D.E. Resasco, Conversion of

1-tetralone over HY zeolite: An indicator of the extent of hydrogen transfer, Appl. Catal. A, 389

(2010) 140-146.


J. Catal., 341 (2016) 136-148.

[4] M. Rueping, A.P. Antonchick, T. Theissmann, A Highly Enantioselective Brønsted Acid

Catalyzed Cascade Reaction: Organocatalytic Transfer Hydrogenation of Quinolines and their

Application in the Synthesis of Alkaloids, Angew. Chem. Int. Ed. , 45 (2006) 3683-3686.






121



A, 102 (1998) 9863-9868.

[8] M.J. Reddish, H.-L. Peng, H. Deng, K.S. Panwar, R. Callender, R.B. Dyer, Direct Evidence

of Catalytic Heterogeneity in Lactate Dehydrogenase by Temperature Jump Infrared

Spectroscopy, J. Phys. Chem. B, 118 (2014) 10854-10862.



[10] F. Lin, Y.-H. Chin, Catalytic Pathways and Kinetic Requirements for Alkanal

Deoxygenation on Solid Tungstosilicic Acid Clusters, ACS Catal., 6 (2016) 6634-6650.

[11] V.B. Kazansky, M.V. Frash, R.A.v. Santen, A quantum-chemical study of hydride transfer

in catalytic transformations of paraffins on zeolites. Pathways through adsorbed nonclassical




[13] A. Miyaji, Y. Iwase, T. Nishitoba, N.Q. Long, K. Motokura, T. Baba, Influence of zeolite

pore structure on product selectivities for protolysis and hydride transfer reactions in the

cracking of n-pentane, Phys.Chem.Chem.Phys., 17 (2015) 5014-5032.

[14] D.A. Simonetti, R.T. Carr, E. Iglesia, Acid strength and solvation effects on methylation,

hydride transfer, and isomerization rates during catalytic homologation of C1 species, J. Catal.,

285 (2012) 19-30.





(2011) 257-268.

[17] F.R. Sarria, V. Blasin-Aubé, J. Saussey, O. Marie, M. Daturi, Trimethylamine as a Probe

Molecule To Differentiate Acid Sites in Y−FAU Zeolite: FTIR Study, The Journal of Physical

Chemistry B, 110 (2006) 13130-13137.

[18] L. Kubelková, J. Čejka, J. Nováková, Surface reactivity of ZSM-5 zeolites in interaction

with ketones at ambient temperature (a FT-i.r. study), Zeolites, 11 (1991) 48-53.

[19] C.D. Chavez Diaz, S. Locatelli, E.E. Gonzo, Acetaldehyde adsorption on HZSM-5 studied

by infrared spectroscopy, Zeolites, 12 (1992) 851-857.

[20] C. Liu, T.J. Evans, L. Cheng, M.R. Nimlos, C. Mukarakate, D.J. Robichaud, R.S. Assary,

L.A. Curtiss, Catalytic Upgrading of Biomass-Derived Compounds via C–C Coupling Reactions:

122

Computational and Experimental Studies of Acetaldehyde and Furan Reactions in HZSM-5, J.

Phys. Chem. C, 119 (2015) 24025-24035.

[21] C.M. Nguyen, M.-F. Reyniers, G.B. Marin, Theoretical study of the adsorption of C1-C4

primary alcohols in H-ZSM-5 Phys. Chem. Chem. Phys., 12 (2010) 9481-9493.

[22] C.M. Nguyen, M.-F. Reyniers, G.B. Marin, Theoretical Study of the Adsorption of the

Butanol Isomers in H-ZSM-5, The Journal of Physical Chemistry C, 115 (2011) 8658-8669.

[23] S. Morin, A. Berreghis, P. Ayrault, N. S. Gnep, M. Guisnet, Dealumination of zeolites Part

VIIIAcidity and catalytic properties of HEMT zeolites dealuminated by steaming, J. Chem. Soc.,

Faraday Trans. , 93 (1997) 3269-3275.



[25] NIST Chemistry WebBook: NIST Standard Reference Database Number 69,

http://webbook.nist.gov/chemistry/. (accessed on Dec 16, 2014).

[26] S. Saito, H. Yamamoto, Design of Acid−Base Catalysis for the Asymmetric Direct Aldol

Reaction, Acc. Chem. Res., 37 (2004) 570-579.

4.6. Appendix

4.6.1. Infrared spectra of H-FAU upon pyridine and butanal adsorption

Figure S4.1a shows the infrared spectra for H-FAU (Si/Al=15) at 373 K, taken after thermal

treatment to 573 K in flowing He (i), followed by either exposure to 0.5 kPa pyridine (ii) or to

0.5 kPa butanal (iii) at 373 K. Butanal adsorption perturbs the hydroxyl stretching bands in the

super- and beta-cages of H-FAU at 3625 cm-1

and 3560 cm-1

[17], respectively, as confirmed by

comparing profiles (i) and (iii) of Figure S4.1a. In fact, these bands disappear as a result of

complete H+ site occupation, as found also during exposure of the sample to pyridine (ii, Fig.

S4.1a). Butanals adsorb on H+ sites as hydrogen-bonded complexes (C3H8C=O

…H

+) [18, 19]

with a characteristic absorption band at 1676 cm-1

[Figure S4.1b, 0.01 kPa C4H8O, H-FAU

(Si/Al=15), 373 K]. The hydrogen-bonded complexes (C3H8C=O…

H+) [18, 19] are chemically

equilibrated with their protonated form (C3H8C+-O-H), because energies between the protonated

and the hydrogen-bonded complexes required for proton transfer remain small. For example, the

http://webbook.nist.gov/chemistry/

123

proton transfer energies for the acetaldehyde [20] and C1-C4 primary alcohols [21, 22] are −6 and

+1~−6 kJ mol−1

, respectively.

Figure S4.1. (a) Infrared spectra for H-FAU in He at 473 K (i), H-FAU exposed to 0.5 kPa

pyridine at 473 K (ii), and H-FAU exposed to 0.5 kPa butanal at 373 K (iii); (b) infrared spectra

for H-FAU exposed to 0.5 kPa butanal at 373 K.

4.6.2. Carbon distribution in the aromatic product fraction during butanal reaction on H-FAU zeolite

Figure S4.2 shows the carbon distribution in the aromatic product fraction during butanal

reaction on H-FAU zeolite at 573 K. The cyclization-dehydration of the alkenal products

(C4tH6t+2O, t=2-5, e.g., 2-ethyl-2, hexenal) formed by the stepwise aldol condensation and

dehydration of butanal produces C8 cyclodienes (C8H12, e.g., dimethyl cyclopentadiene) or C12,

C16, and C20 aromatic species (C4tH6t, t=3-5, e.g., triethylbenzene) [3, 10]. The stepwise

dehydrogenation of these primary cyclization products can further increase their extents of

unsaturation [3]. Meanwhile, the transalkylation of these aromatic species leads to the formation

of aromatic products ranging from C7 to C19 [3], as shown in Figure S4.2.

4000 3800 36000.0

0.3

0.6

0.9

1.2

1.5

1.8

Butanal-FAU

Pyridine-FAU

3560

(OH

) beta

3625

(OH

) super

No

rma

lize

d a

bso

rba

nce

(a

.u.)

Wavenumber (cm-1)

3740

(Si-

OH

)

Pristine FAU

1700 1600 1500 1400

-1

0

1

2

Ab

so

rba

nce

(a

.u.)

Wavenumber (cm-1)

14

69

δas(C

H3)

14

90

ν(C

=C

) aro

m.

14

08

δ(C

H)

16

00

ν(C

=C

) aro

m.

16

76

ν(C

=O

)

(b)(a)

(i)

(ii)

(iii)

124

Figure S 4.2. Carbon distributions of aromatic fraction produced in butanal reactions on H-FAU

at 573 K (time-on-stream 125 min). The distributions include aromatic molecules that do not lose

any H ( ) or lose 2 ( ), 4 ( ), or 6 ( ) hydrogen atoms in dehydrogenation reactions.

Examples of the C12 aromatic products are given. These results have been published elsewhere

[3].

4.6.3. Determination of aldehyde transfer hydrogenation rate by co-feed H-donors

During the reaction of aldehyde (CnH2nO, n=3-6) on H-FAU zeolite, alkene (CnH2n) is formed

via the transfer hydrogenation of aldehyde (CnH2nO) by the aromatic products as H-donors,

which are produced by the intermolecular pathway of aldehyde condensation and the sequential

cyclization [9]. In the mixture of aldehyde and co-feed H-donor (denoted as RDH2), both the

aromatic products and RDH2 can act as H-donors, contributing to the formation of alkene (CnH2n).

The rates for aldehyde (CnH2nO) transfer hydrogenation in the presence of external H-donors

(RDH2), 𝑟TH,C𝑛H2𝑛O−RDH2, are isolated by subtracting the site-time-yields of alkene (CnH2n) in

CnH2nO feed (𝑟Intra,C𝑛H2𝑛O ) from those in CnH2nO-RDH2 feed mixture (𝑟Intra,C𝑛H2𝑛O−RDH2),

measured on H-FAU (Si/Al=15) zeolites at 573 K:

𝑟TH,C𝑛H2𝑛O−RDH2= 𝑟Intra,C𝑛H2𝑛O−RDH2

− 𝑟Intra,C𝑛H2𝑛O (S4.1)

0

5

10

15

C16H

20

C16H

22

C16H

24

C15H

16

C15H

18

C15H

20

C15H

22

C15H

24

C7H

8

C8H

10

C9H

12

C11H

10

C11H

12

C11H

14

C11H

16

C10H

14

C10H

10

C10H

12

C10H

8

Ca

rbo

n d

istr

ibu

tio

n in

aro

ma

tics (

%)

Aromatic products

C12H

12

C12H

14

C12H

16

C12H

18

C13H

14

C13H

16

C13H

18

C13H

20

C14H

22

C14H

20

C14H

18

C14H

16

C17H

22

C17H

24

C17H

26

C18H

24

C18H

26

C18H

28

C19H

30

C19H

28

C19H

26


125

4.6.4. Kinetic relevance of hydrogen transfer step in aldehyde transfer hydrogenation on Brønsted acid sites

Scheme 4.3a in the main manuscript depicts the mechanistic steps for the aldehyde transfer

hydrogenation. First, the hydrogen transfer step (Step 1, Scheme 4.3a), the H-donor (denoted as

RDH2, e.g., aromatic products) donates a hydride ion onto the protonated aldehyde (CnH2nOH+),

forming an alcohol (CnH2n+1OH); at the same time, H-donor is converted into carbenium ion

(RDH+) which then donates a proton to regenerate the H

+ site (Step 2b, Scheme 4.3a); the

sequential dehydration of the alcohol (Step 3, Scheme 4.3a) forms an alkene (CnH2n) and

completes the catalytic cycle. I have reported in our previous study[3, 10] that the rates of alkene

formation is first order on the pressure of H-donors (e.g., aromatic products and co-feed tetralin);

in addition, the rate of alkene formation via alcohol dehydration is a least one order of magnitude

higher than that via aldehyde transfer hydrogenation, under identical conditions (temperature and

partial pressure) [3, 9]. These results indicate that the hydrogen transfer step, instead of the

alcohol dehydration step, is the kinetically relevant step for the aldehyde transfer hydrogenation.

4.6.5. H+ site coverage by butanal and its derivatives in infrared

spectroscopic study

Figure S4.3a shows the differential infrared spectra for H-FAU during butanal adsorption and

desorption by subtracting the spectrum of pristine H-FAU. The bands at 3626 cm-1

and 3560 cm-1

are ascribed to the stretching vibrations of H+ sites in the supercages and the beta cages,

respectively; the bands 3740 cm-1

and 3600 cm-1

are ascribed to the non-acidic silanol groups

(Si-O-H) and the weak-acidic hydroxyl groups on extra-framework alumina (Al-O-H) [23],

respectively. During the initial 270 s, butanal (10 Pa) was feed in the cell, butanal and its

derivatives (e.g., 2-ethyl-2-hexenal and aromatics) adsorbed on the H+ sites, leading to the

negative peaks at 3626 and 3560 cm-1

. The heavy condensation products (e.g., 2-ethyl-2-hexenal

and aromatics) could also adsorb on the Si-O-H and Al-O-H, resulting in the negative peaks at

126

3740 and 3600 cm-1

, respectively. After 270 s when the butanal feed was stopped and the sample

was purged with He, butanal desorption from the H+ sites led to the recovery of the bands at

3626 and 3560 cm-1

. However, Al-O-H (3600 cm-1

) and Si-O-H (3740 cm-1

) did not recover,

because the heavy products, which were formed upon the condensation reaction between the

protonated butanal and gaseous butanal and the sequential cyclization, could migrate and stay on

Al-O-H and Si-O-H groups. The total H+ site coverage (𝜃Total) was determined by Equation S4.2:

𝜃Total =𝐴𝑣(OH),pristine sample−𝐴𝑣(OH),butanal adsorbed sample

𝐴𝑣(OH),pristine sample (S4.2)

where 𝐴𝑣(OH),pristine sample and 𝐴𝑣(OH),butanal adsorbed sample are the intensities of the H+ site

stretching bands (including bands at 3626 cm-1

and 3560 cm-1

) of the infrared spectra of pristine

H-FAU and butanal adsorbed H-FAU, respectively. Figure S4.3b shows the total H+ site coverage

(𝜃Total) during butanal adsorption and desorption as a function of time, showing that the H+ site

coverage decreased upon helium purging starting at 270 s. However, the H+ sites could not be

completely recovered, because some heavy condensation products (e.g., 2-ethyl-2-hexenal and

aromatics) could not be desorbed at the low temperature of 373 K.

Figure S4.3. (a) Differential infrared spectra of H-FAU during butanal adsorption and desorption

by subtracting the spectrum of pristine H-FAU; (b) the total H+ site coverage during butanal

127

adsorption and desorption on H-FAU as a function of time (373 K, butanal pressure 10 Pa,

helium flow rate 0.83 cm3 s

−1).

Table S4.1 summarizes the initial total H+ site coverage by butanal and its derivatives (𝜃Total)

and H+ site coverage by butanal (𝜃C4H8O,0) on H-FAU before feeding H-donors in each in-situ

FTIR experiment. 𝜃Total was determined using the H+ site bands at 3626 and 3560 cm

-1

according to Equation S4.2, whereas 𝜃C4H8O,0 was calculated based on the band for protonated

carbonyl group at 1675 cm-1

(Figs. 4.2a-4.2b and Figs. S4.5a-S4.5c) and the details of the

calculation are in Section 4.6.6.

Table S4.1. Initial total H+ site coverage by butanal and its derivatives ( Total ) and H

+ site

coverage by butanal (4 8C H O,0 ) on H-FAU before feeding H-donors.

H-donors Total (ML)

4 8C H O,0 (ML)

No H-donor 0.12±0.05 0.09

Cyclohexane 0.13±0.05 0.08

Cyclohexene 0.11±0.05 0.08

Tetralin 0.11±0.05 0.09

Cyclohexadiene 0.16±0.05 0.12

4.6.6. Estimation of the H+ site coverage by carbonyl group on H-

FAU

The H+ sites coverage by butanal was calculated via the measured aldehyde coverage on H-MFI

to deduce on H-FAU. Figure S4.4a shows the infrared spectra of butanal adsorption on H-MFI

zeolite (Si/Al=40) at 308 K in 0.01 Pa butanal. The H+ site occupation by the adsorbed butanal is

indicated by the decrease in the bridging hydroxyl (SiOHAl, the H+ site) stretching band [v(OH)]

128

at 3610 cm−1

and the concomitant increase in the carbonyl stretching band [v(C=O)] at 1670

cm−1

. Because the H+ sites are occupied solely by the butanal monomer, the decrease of v(OH)

band area (−∆𝐴𝑣(OH),MFI) is proportional to the increase of the v(C=O) band area (∆𝐴𝑣(C=O),MFI),

as shown in Figure S4.4b.

The coverage of H+ sites by butanal (𝜃C4H8O) is determined based on the area ratio of the changes

in v(OH) band caused by C=O occupation (−∆𝐴𝑣(OH),FAU) to the v(OH) band of pristine FAU

sample (𝐴𝑣(OH),FAU,0), according to Equation S4.3:

𝜃C4H8O =−∆𝐴𝑣(OH),FAU

𝐴𝑣(OH),FAU,0 (S4.3)

Because a portion of H+ sites are occupied by heavy products (e.g., aromatics), −∆𝐴𝑣(OH),FAU

(corresponding to the C=O occupation) could not be directly measured on the IR spectra, and its

value needs to be derived based on the changes in the band of the C=O group (∆𝐴𝑣(C=O),FAU).The

extinction coefficients (𝜀𝑣(OH),MFI and 𝜀𝑣(OH),FAU) for the bridging hydroxyl group (SiOHAl, H+

site) on FAU and MFI zeolite are evaluated to be 3.7 and 3.1, respectively, according to previous

report [24]. Therefore, the changes in the amount of H+ sites on these two zeolites (∆[OH]MFI

and ∆[OH]FAU, respectively) are correlated to the changes in their IR band (∆𝐴𝑣(OH),MFI and

∆𝐴𝑣(OH),FAU, respectively) according to Equation S4.4:

∆[OH]MFI

∆[OH]FAU=

∆𝐴𝑣(OH),MFI/𝜀𝑣(OH),MFI

∆𝐴𝑣(OH),FAU/𝜀𝑣(OH),FAU=

∆𝐴𝑣(OH),MFI/3.7

∆𝐴𝑣(OH),FAU/3.1 (S4.4)

Assuming the extinction coefficient for the C=O stretching band [v(C=O)] is independent on the

catalyst identities (𝜀𝑣(C=O),MFI =𝜀𝑣(C=O),FAU) , the changes in the amount of C=O on these two

samples (∆[C=O]MFI and ∆[C=O]FAU, respectively) are correlated to the changes in their IR band

(∆𝐴𝑣(C=O),MFI and ∆𝐴𝑣(C=O),FAU, respectively) according to Equation S4.5:

∆[C=O]MFI

∆[C=O]FAU=

∆𝐴𝑣(C=O),MFI/𝜀𝑣(C=O),MFI

∆𝐴𝑣(C=O),FAU/𝜀𝑣(C=O),FAU=

∆𝐴𝑣(C=O),MFI

∆𝐴𝑣(C=O),FAU (S4.5)

Assuming each adsorbed C=O group only occupies one H+ site, I have Equation S4.6:

∆[OH]MFI

∆[C=O]MFI=

∆[OH]FAU

∆[C=O]FAU= 1 (S4.6)

129

By combining Equations S4.4-S4.6, I obtain the correlation among the changes in IR bands of

the H+ site and C=O group on H-MFI and H-FAU in Equation S4.7:

−∆𝐴𝑣(OH),MFI/3.7

∆𝐴𝑣(C=O),MFI/1=

−∆𝐴𝑣(OH),FAU/3.1

∆𝐴𝑣(C=O),FAU/1 (S4.7)

which is re-organized to Equation S4.8:

−∆𝐴𝑣(OH),FAU =3.1

3.7∙

−∆𝐴𝑣(OH),MFI

∆𝐴𝑣(C=O),MFI∙ ∆𝐴𝑣(C=O),FAU (S4.8)

The value of −∆𝐴𝑣(OH),MFI(∆𝐴𝑣(C=O),MFI) is 0.1945, according to Figure S4.3b. Therefore,

Equation S4.8 is rewritten as:

−∆𝐴𝑣(OH),FAU = 0.163 ∙ ∆𝐴𝑣(C=O),FAU (S4.9)

By substituting Equation S4.9 into Equation S4.3, I obtain the correlation between the H+ site

coverage by butanal (𝜃C4H8O) and the intensity of C=O band (𝐴𝑣(C=O),FAU), as shown in Equation

S4.10:

𝜃C4H8O = 0.163 ∙∆𝐴𝑣(C=O),FAU

𝐴𝑣(OH),FAU,0= 0.163 ∙

𝐴𝑣(C=O),FAU−0

𝐴𝑣(OH),FAU,0= 0.163 ∙

𝐴𝑣(C=O),FAU

𝐴𝑣(OH),FAU,0 (S4.10)

Figure S4.4. (a) Infrared spectra of butanal adsorption on H-MFI zeolite (Si/Al=40) at 308 K in

0.01 Pa butanal; (b) correlation between the decrease of band area for the H+ site (−∆Av(OH), MFI,

1800 1600 1400

0.0

0.2

0.4

0.6

0.8

Ab

so

rba

nce

(a

.u.)

Wavenumber (cm-1)

3700 3600 3500

0.0

0.1

0.2

0.3

0.4

0.5

Ab

so

rba

nce

(a

.u.)

Wavenumber (cm-1)

0 10 20 30 400

2

4

6

8

A

v(O

H),

MF

I

Av(C=O),MFI

(a) (b)

130

stretching vibration at 3610 cm−1

) and the increase of the band area for the adsorbed carbonyl

group (∆Av(C=O), stretching vibration at 1670 cm−1

) during butanal adsorption on H-MFI zeolite

(Si/Al=40) at 308 K in 0.01 Pa butanal.

Figure 4.2b (in Sec.4.3.1) and Figures S4.5a-S4.5c show the time-resolved infrared spectra of H-

FAU zeolite upon butanal adsorption followed by butanal desorption in the presence of various

H-donors (RDH2=tetralin, cyclohexadiene, cyclohexene, and cyclohexane, respectively) at 373 K.

The band for protonated carbonyl group v(C=O) at 1676 cm−1

increased during butanal feeding

as butanal adsorbed on H+ sites (gray lines, Fig. 4.2b and Figs. S4.5a-S4.5c), and then began to

decrease after stopping butanal feed due to butanal desorption and transfer hydrogenation (black

lines, Fig. 4.2b and Figs. S4.5a-S4.5c). The rates for v(C=O) band decreasing varied with the

identities of the introduced H-donors, because of their different efficiencies in butanal transfer

hydrogenation. The H+ site coverages by butanal ( 𝜃C4H8O ) during these experiment were

determined based on the intensities of the v(C=O) band at 1676 cm−1

according to Equation

S4.10, and plotted in Figure 4.2c in the main manuscript.

Figure S4.5. Time-resolved infrared spectra of H-FAU zeolite upon butanal adsorption (10 Pa

butanal, gray lines) followed by purging in (a) cyclohexadiene (10 Pa), (b) cyclohexene (11 Pa),

and (c) cyclohexane (12 Pa) at 373 K.

131

4.6.7. Comparison of rates for intermolecular C=C bond formation with and without H-donor incorporation

The incorporation of H-donors (RDH2) in the butanal (C4H8O) reaction does not influence the

rates for the intermolecular C=C bond formation (RDH2=cyclohexadiene, tetralin, cyclohexene,

3-methyl-1-pentene, and cyclohexane), as indicated by the rate ratios ( 𝑟Inter,C4H8O−RDH2/

𝑟Inter,C4H8O) of almost unity in Table S4.2 (𝑟Inter,C4H8O−RDH2 and 𝑟Inter,C4H8O denote the rates for

intermolecular C=C bond formation in C4H8O-RDH2 feed mixture and in C4H8O feed,

respectively).

Table S4.2. Rate ratios for the pathway of intermolecular C=C bond formation in C4H8O-RDH2

feed mixture (D4 8 2,C H O R HInter -r ) to that in C4H8O feed (

4 8,C H OInterr ) on H-FAU at 573 K

RDH2 4 8 D 2,C H O R HInter -r /

4 8,C H OInterr

Cyclohexadiene 1.07

Tetralin 0.97

Cyclohexene 1.02

3-methyl-1-pentene 0.99

Cyclohexane 1.02

4.6.8. Estimation of hydride ion affinities for protonated aldehydes and carbenium ions of H-donors

The hydride ion affinities (HIA) for the carbenium ions (RDH+) of the selected hydrocarbons

(RDH2, e.g., tetralin C10H12) and the protonated aldehydes (CnH2nOH+, n=3, 4, 5, 6) (denoted as

D+R H

HIA and +2C H OHn n

HIA , respectively) are defined as the heats of hydride ion addition reactions

as shown in Equations S4.11 and S4.12, respectively:

RDH++ H

−→RDH2 (S4.11)

CnH2nOH++ H

−→CnH2n+1OH (S4.12)

132

D+R H

HIA and +2C H OHn n

HIA were estimated based on the thermochemical cycles shown in Scheme

S4.1, which use tetralin (Scheme S4.1a) and alkanal (Scheme S4.1b) to illustrate the method.

The values of D

+R HHIA (or +

2C H OHn nHIA ) were calculated according to Equations S4.13-S4.16,

using proton affinities ( PA ), heats of formation ( f H ) of the hydrogen donors RDH2 (or the

aldehyde hydrogenation products CnH2n+1OH), the diatomic hydrogen H2, and the

dehydrogenation products RD (or the aldehydes, CnH2nO), and heat of reaction for the reaction

between a proton and a hydride (H++H

−→H2, r H ionH =−1675.3 kJ mol

-1 [25]). The estimated

values are listed in Table S4.3.


the carbenium ion (RDH+) of a hydrocarbon (RDH2, taking RDH2=tetralin as an example) and (b)

the protonated aldehyde (CnH2nOH+) [

DRPA and 2C H On n

PA are the proton affinities of

hydrocarbon RD and aldehyde CnH2nO, respectively; r H ionH (−1675.3 kJ mol-1

) [25] is the heat

of reaction for H++H

- → H2;

Dr Hydro, RH and 2, C H Or Hydro n n

H are the heats of reaction for

hydrogenation reactions: RD+H2→ RDH2 and CnH2nO+ H2→CnH2n+1OH, respectively].

D+R H

HIA =DRPA + r H ionH +

Dr Hydro, RH (S4.13)

and

Dr Hydro, RH =D 2R Hf H −

DRf H −2Hf H (S4.14)

(RDH2)

(RD)

(RDH+)

(RD)

(CnH2n+1OH)(CnH2nOH+)

(CnH2nO) (CnH2nO)

(a) (b)

133

+2C H OHn n

HIA =2C H On n

PA + r H ionH + , 2r Hydro C H On nH (S4.15)

and

, 2r Hydro C H On nH =

2 1C H OHf n nH

−

2C H Of n nH −

2Hf H (S4.16)

Table S4.3. Hydride ion affinities (HIA) for protonated aldehydes (CnH2nOH+; n=3-6) and the

carbenium ions of the H-donors (RDH+).

Carbenium ion

(RDH+) or protonated

aldehyde

(CnH2nOH+)

H-donor (RDH2) or

aldehyde (CnH2nO)

DRPA or

2C H On nPA

(kJ mol−1

)a

DRf H or

2C H Of n nH

(kJ mol−1

)b

D 2R Hf H or

2 1C H OHf n nH

(kJ mol−1

)b

D+R H

HIA or

+2C H OHn n

HIA

(kJ mol−1

)

C3H6OH+ Propanal 786.0 -188.7 -256 956.6

C4H8OH+ Butanal 792.7 -211.8 -277 947.8

C5H10OH+ Pentanal 796.6 -232.4 -298 944.3

C6H12OH+ Hexanal 801.6 -248.6 -316 941.1

C10H11+ Tetralin 842.0 [26] 130.8 30.0 934.1

C6H7+ Cyclohexadiene 746.4 82.9 104.6 907.2

C6H9+ Cyclohexene 837 104.58 -3.32 946.2

C6H11+ Cyclohexane 784 -3.32 124.6 1019.2

C6H11+ 3-methyl-1-pentene 852.3 75.3 -50.1 948.4

C2H5+ Ethane 680.5 52.5 -84 1131.3

C3H7+ Propane 751.6 20.41 -104.7 1048.8

i-C4H9+ i-Butane 802.1 -17.9 -134.2 989.5

aProton affinities (PA) were obtained from ref [12, 25]. bHeats of formation were obtained from the database [25].

134

Chapter 5 Catalytic Pathways and Kinetic Requirements for Alkanal

Deoxygenation on Solid Tungstosilicic Acid Clusters

Abstract

Kinetic measurements and acid site titrations were carried out to interrogate the reaction network,

probe the mechanism of several concomitant catalytic cycles, and explain their connection

during deoxygenation of light alkanals (CnH2nO, n=3-6) on tungstosilicic acid clusters

(H4SiW12O40) that leads to hydrocarbons (e.g., light alkenes, dienes, and larger aromatics) and

larger oxygenates (e.g., alkenals). The three primary pathways are: (1) intermolecular C=C bond

formation, which couples two alkanal molecules in aldol condensation reactions followed by

rapid dehydration, forming a larger alkenal (C2nH4n-2O), (2) intramolecular C=C bond formation,

which converts an alkanal directly to a n-alkene (CnH2n), by accepting a hydride ion from H

donor and ejecting a H2O molecule, and (3) isomerization-dehydration, which involves self-

isomerization of an alkanal to form an allylic alcohol and then rapidly dehydrate to produce a n-

diene (CnH2n-2). The initial intermolecular C=C bond formation is followed by a series of

sequential intermolecular C=C bond formation steps, during each of these steps an additional

alkanal unit is added onto the carbon chain to evolve a larger alkenal (C3nH6n-4O and C4nH8n-6O),

which upon its cyclization-dehydration reaction forms hydrocarbons (CtnH2tn-2t, t=2-4, including

cycloalkadienes or aromatics). The inter- and intramolecular C=C bond formation cycles are

catalytically coupled through intermolecular H-transfer events, whereas the intermolecular C=C

bond formation and isomerization-dehydration pathways share a co-adsorbed alkanal-alkenol

pair as the common reaction intermediate. The carbon number of alkanals determines their

hydride ion affinities, the stabilities of their enol tautomers, and the extent of van der Waals

interactions with the tungstosilicic acid clusters; these factors influence the stabilities of the

transition states or the abundances of reaction intermediates in the kinetically-relevant steps and

in turn the reactivities and selectivities of the various cycles.

135

5.1. Introduction

Fast pyrolysis of lignocellulosic biomass produces light oxygenates with less than or equal to six

carbon atoms [1, 2]. Contained within the light oxygenate fraction are alkanals, such as

hydroxyacetaldehyde and furfural, which account for ~20 wt.% of the organic fraction [2, 3].

These alkanals react on solid Brønsted acid catalysts (e.g., H-ZSM-5 [4-7], H-MOR [6], H-

FAU[6, 7] zeolites) via a series of aldol condensation and dehydration reactions, through which

they augment their size by creating intermolecular carbon-carbon linkages. The condensation

reactions may occur multiple times to further augment the carbon chain until the eventual

intramolecular carbon-carbon bond formation, followed by dehydration, dehydrogenation, and

transalkylation to evolve diverse aromatics. As an example, deoxygenation of propanal (C3H6O)

on H-ZSM-5 zeolites at 673 K leads predominantly to C6-C10+ aromatics with carbon

selectivities between 42 % and 53 % [8, 9].

The mechanism for the initial aldol condensation on solid acid catalysts (H-MFI[10, 11] and H-

Y[12]) has been well established, but few studies have addressed the sequential reactions that

lead to the formation of larger olefinic or aromatic products. Propanal reactions on H-ZSM-5

zeolites involve self-condensation and dehydration steps that form the dimeric species (2-methyl-

2-pentenal, C6H10O), which undergo sequential cross condensation with another propanal to

produce trimeric species (2,4-dimethyl-2,4-heptadienal, C9H14O), before their ring-closure and

dehydration to evolve C9 aromatics [8, 9]. These C9 aromatics then undergo secondary

transalkylation steps that shuffle their alkyl groups via carbenium ion transfer[13] and result in

C6-C9+ aromatics [8, 9].

Other reactions occur concurrently with the intermolecular carbon-carbon bond formation and

ring closure reactions. Alkanals (CnH2nO) may remove their oxygen via a direct dehydration

route, which forms the corresponding dienes (CnH2n-2) [14-17]. In fact, previous studies have

shown that 2-methylbutanal dehydration on borosilicate zeolite[14] or aluminum phosphate

(AlPO4)[16, 17] leads to isoprene [14, 16, 17], whereas 2-methylpentanal dehydration on

aluminosilicate zeolite (H-Y) leads to 2-methylpenta-1,3-diene[14], as viable routes for

synthesizing polymer precursors. These alkanal dehydration reactions were proposed[17] to

occur via a common allylic alcohol intermediate: 2-methylbutanal reactions catalyzed by BPO4

and AlPO4 catalysts (598-673 K) form isoprene and methyl isopropyl ketone; at the similar

136

conditions, both 2-methylbutanal and methyl isopropyl ketone reactions give similar yields to

isoprene on AlPO4 (54 % vs. 49 % at 673 K). Therefore, these reactions must involve a common

allylic intermediate for the interconversion between isoprene and methyl isopropyl ketone [17].

In addition, 2-methyl-2-buten-1-ol reaction on BPO4 (383 K) forms 2-methylbutanal, methyl

isopropyl ketone, and isoprene with selectivities of 11 %, 46 %, and 43 %, respectively. These

allylic alcohols, alkanals, ketones, and isoprenes can interconvert with the allylic alcohol as the

intermediate [17]. During alkanal dehydration, the formation of allylic alcohol is likely the initial

kinetically relevant step, because the 2-methyl-2-buten-1-ol remains undetected during 2-

methylbutanal dehydration (BPO4 and AlPO4) at 598-673 K [17].

A separate reaction for alkene formation from alkanal may also occur, as reported previously for

alkanal reactions on H-ZSM-5 zeolite [8, 11]. Propanal reactions on H-ZSM-5 zeolite at 673 K

produce a significant amount of C1-C3 light gases (43-53% carbon selectivities) and

predominantly propene [8]. In fact, reactions of CnH2nO alkanal (n=3-5) on H-ZSM-5 zeolites

produce almost exclusively CnH2n alkenes [CnH2n/(1

21C H

t n

t tt n

)=0.93, 0.95, and 0.89 for n=3,

4, and 5, respectively, at 473 K] within the alkene product fraction [11]. The alkene formation

likely occurs via a direct hydrogen transfer step, during which a protonated alkanal accepts a

hydride ion, followed by dehydration and desorption as alkene, leaving its carbon backbone

intact [18].

Several catalytic routes occur concomitantly, which result in larger oxygenates, alkenes,

aromatics, as well as light alkenes and dienes during alkanal deoxygenation on solid Brønsted

acid catalysts. Their individual rates, kinetic requirements, and the kinetic connection between

these pathways have, however, remained largely unresolved. The ambiguity of the catalytic

pathways and the associated mechanism are caused, in large part, by the complexity of the

reaction systems, which appear to involve condensation of two alkanals, dehydration of a single

alkanal, shuffling of H atoms from products to reactants, and various secondary ring closure and

transalkylation reactions. Probing these inherently complex pathways on catalysts containing

diverse site structures further complicates the rate data interpretation, because rates of these steps

are expected to vary with the site structures and their thermodynamic properties.

Here, I probe the catalytic pathways of alkanal deoxygenation with kinetic and chemical titration

strategies, after isolating the kinetic contributions of acid site and site environment. I focus on

137

the deoxygenation chemistry of straight chain alkanals with three to six carbon atoms (CnH2nO,

n=3-6), carried out on tungstosilicic acid clusters (H4SiW12O40) with well-defined structures.

Such clusters contain isolated H+ sites without the local molecular confinement typically found

in microporous crystalline materials. Through quantitative kinetic studies, I probe the reaction

pathways by systematically examining the primary and secondary reactions and also by

decoupling the rate contributions from the various catalytic routes. Specifically, I establish the

kinetic correlation between the three primary pathways during C3-C6 alkanal (CnH2nO)

deoxygenation on Brønsted acid sites of H4SiW12O40 clusters that lead to larger alkenals (C2nH4n-

2O) through bimolecular C=C bond formation, light alkenes (CnH2n) via H-transfer and

dehydration, and dienes (CnH2n-2) from direct dehydration reactions. These rates and selectivities

on tungstosilicic acid clusters differ from those on microporous crystalline materials (e.g., H-

MFI[11] and H-FAU[18]); specifically, the tungstosilicic acid clusters exhibit much higher

selectivities towards alkanal coupling than H-transfer and direct alkanal dehydration reactions

and less extent of cyclization and transalkylation reactions, because of the lack of local H+ site

confinements and the different extents of van der Waals interaction compared to zeolites. Our

approach provides simple explanations to the apparent complex reaction system and correlates

thermochemical properties (e.g., hydride ion affinities and heats of adsorption) to rates and

selectivities during deoxygenation reaction.

5.2. Experimental

5.2.1. Preparation and characterizations of H4SiW12O40 clusters dispersed on SiO2 support

H4SiW12O40/SiO2 catalysts (loading amount=0.075 4 12 40 2

1H SiW O SiO

mmol g ) were prepared by

incipient wetness impregnation method. SiO2 support (GRACE chromatographic grade, Code

1000188421, surface area=330 m2 g

−1, particle size<75 μm, pore volume=1.2 cm

3 g

−1) was

heated in air (Linde, zero grade) at 0.17 K s−1

to 773 K and then maintained at 773 K for 5 h. The

treated SiO2 support was impregnated with a solution at a liquid-to-solid ratio of 1.2 cm3 gSiO2

−1,

prepared by dissolving H4SiW12O40 as received (Sigma Aldrich, reagent grade, CAS #12027-43-

9) in ethanol (Sigma-Aldrich, >99.5 %, anhydrous). The sample was then held in a closed vial

138

for 24 h and then treated in flowing dry air [Linde, zero grade, 0.1 cm3 (gcat.

s)

−1] at 0.017 K s

−1

to 323 K and maintained at 323 K for 24 h.

The ratio of Brønsted-to-Lewis acid sites on H4SiW12O40/SiO2 catalysts was determined by the

infrared spectroscopic study of pyridine adsorption at 473 K. The Brønsted-to-Lewis site ratio

was found to be 14.7, as shown in Appendix (Sec. 5.6.1) [19]. The total acid site densities

(including Brønsted and Lewis sites) were determined by isothermal chemical titration with

pyridine followed by temperature programmed desorption (TPD) in flowing He. Catalyst

powders (150 mg) were loaded into a microcatalytic quartz reactor (9.5 mm inner diameter),

supported on a coarse quartz frit. The catalyst powders were treated in-situ under flowing He

(Linde, Grade 5.0, 0.83 cm3 s

−1) at a constant heating rate of 0.083 K s

−1 to 473 K. As the reactor

temperature reached and maintained isothermally at 473 K, pyridine (Sigma Aldrich, >99.9 %,

CAS#110-86-1) was introduced at 3.42×10−8

mol s−1

through a gas tight syringe (SGE, Model

006230, 0.25 cm3) into a vaporization zone maintained at 391 K and located at the upstream of

the reactor, within which pyridine was evaporated and mixed with a flowing He stream (Linde,

Grade 5.0, 0.83 cm3 s

−1). The amount of pyridine in the effluent stream was quantified using a

flame ionization detector (FID) in a gas chromatograph (Agilent, 7890A). Pyridine adsorption

was completed when the molar flow rate of pyridine in the effluent stream became identical to

that of the feed stream, at which point the isothermal chemical titration step was completed. The

reactor was subsequently purged in flowing He (Linde, Grade 5.0, 0.83 cm3 s

−1) at 473 K for 30

min. The He flow rate was then adjusted to 0.17 cm3 s

−1 and the temperature was increased

linearly from 473 K to 923 K at 0.033 K s−1

. The amount of pyridine desorbed into the effluent

stream as a function of time (which was also related to the temperature) was quantified using the

FID detector. The total acid site densities were determined based on the pyridine uptakes during

the chemical titration step as well as that of pyridine desorbed during the TPD, by assuming a

pyridine-to-acid site molar ratio of unity. Both methods gave consistent results (0.169±0.006

1cat.acid site

mmol g ), thus the Brønsted site density is 0.159±0.006 +

1cat.H

mmol g based on the

Brønsted-to-Lewis site ratio determined by the infrared spectra of pyridine adsorption. The

turnover rates of alkanal reactions reported in this work were calculated based on the initial H+

site density on the fresh H4SiW12O40/SiO2 catalysts.

139

The H+ site titration with alkanal (CnH2nO, n=3-6) was performed using a procedure similar with

the pyridine titration. 50 mg of catalyst powders were loaded in the microcatalytic quartz reactor.

The samples were treated under flowing He (Linde, Grade 5.0, 0.83 cm3 s

−1) by heating to 473 K

at 0.083 K s−1

, held for 0.5 h at 473 K, and then cooled to 348 K. The alkanal [propanal (Sigma

Aldrich, Kosher grade, ≥97 %, CAS #123-38-6), butanal (Sigma Aldrich, puriss grade, ≥99 %,

CAS #123-72-8), pentanal (Sigma Aldrich, 97%, CAS #110-63-3), or hexanal (Sigma Aldrich,

≥98 %, CAS#66-25-1)] was introduced at 1.7×10−8

mol s−1

through a gas tight syringe (SGE,

Model 006230, 0.25 cm3) into a vaporization zone, which was maintained at the boiling point of

the alkanal, within which the alkanal was evaporated and mixed with a flowing He stream (Linde,

Grade 5.0, 0.83 cm3 s

−1). The amount of alkanal in the effluent stream was quantified using a

flame ionization detector (FID) in a gas chromatograph (Agilent, 7890A). Alkanal adsorption

was completed when the molar flow rate of alkanal in the effluent stream became identical to

that of the feed stream.

5.2.2 Rate and selectivity assessments for alkanal deoxygenation on H4SiW12O40 polyoxometalate clusters

Reactions of alkanals (CnH2nO, n=3-6) or 2,4-heptadienal (C7H10O) on H4SiW12O40/SiO2

catalysts were carried out in a fixed bed microcatalytic quartz reactor (9.5 mm inner diameter)

with plug-flow fluid dynamics at 573 K. The reactor was contained within a resistively heated

furnace with its temperature controlled by a digital feedback controller (Omega, CN3251). Inside

the quartz reactor, catalyst powders (25 or 50 mg) were supported on a coarse quartz frit and the

bed temperature was recorded using a K-type thermocouple placed at the center (in both the axial

and radial directions) of the catalyst bed. Catalysts were treated in-situ under flowing He [Linde,

Grade 5.0, 4.16-33.3 cm3 (gcat.

s)

-1], by heating at 0.167 K s

-1 to the reaction temperature (573 K

or 623 K) prior to rate and selectivity measurements. Propanal (Sigma Aldrich, Kosher grade,

≥97 %, CAS #123-38-6), butanal (Sigma Aldrich, puriss grade, ≥99 %, CAS #123-72-8),

pentanal (Sigma Aldrich, 97%, CAS #110-63-3), hexanal (Sigma Aldrich, ≥98 %, CAS#66-25-1),

or 2,4-heptadienal (Sigma Aldrich, ≥90 %, CAS# 4313-03-5) was introduced into a vaporization

zone located at the upstream of the reactor through a gas-tight syringe (Hamilton, Gastight 1105,

5 mL, or SGE, Model 006230, 0.25 cm3), mounted on a syringe infusion pump (KD Scientific,

LEGATO 100). In the vaporization zone, the reactant was evaporated and mixed with a flowing

He stream [Linde, Grade 5.0, 4.16-33.3 cm3(gcat.

s)

−1]. The partial pressure of reactants was

140

maintained at a constant value between 1.1 kPa to 10 kPa by controlling the liquid infusion rate

of the syringe infusion pump. The mixture was fed to the reactor via heated transfer lines held at

473 K. The reactor effluent stream was kept above 473 K and quantified with an on-line gas

chromatograph (Agilent, 7890A) and mass spectrometer (Agilent, 5975C) equipped with two

capillary columns of (i) Agilent HP-5MS (190091S-433, 30 m, 0.25 mm ID, 0.25 μm film)

connected to a thermal conductivity detector (TCD) and a flame ionization detector (FID) in

series and (ii) HP-5 (19091J-413, 30 m, 0.32 mm ID, 0.25 μm film) connected to the mass

spectrometer. These two capillary columns separated the effluent species in the same order and

with very similar retention times. After chromatographic separation, each peak which

corresponds to a chemical species was identified by examining its associated mass spectrum and

then matching the mass spectrum to the NIST/EPA/NIH mass spectral library. Using this method,

peaks corresponding to hydrocarbons (olefins, aromatics, dienes, etc.) and oxygenates (alkenals,

alkenones, etc.) were identified. The concentrations of these species were further quantified

based on their individual FID signal intensity and FID response factor (determined according to

the method established in the literature[20]). The CO and CO2, which could not be detected by

FID, were quantified based on their relative mass spectrum signal intensities in comparison with

the hydrocarbon species (e.g., C3-C6 alkenes).


5.3.1. Catalytic pathways of alkanal deoxygenation on H4SiW12O40 tungstosilicic acid dispersed on high surface area silica substrates

Reactions of straight chain alkanals (CnH2nO, n=3-6) on solid Brønsted acid sites at moderate

temperatures (473-673 K) and the ambient pressure form larger alkenals and their isomers

(CtnH2tn-2t+2O, n=3-6, t=2-4), as well as hydrocarbons including aromatics (CtnH2tn-2t, n=3-6, t=3-

4), cycloalkadienes (CtnH2tn-2t, n=3-6, t=2), cycloalkenes (CtnH2tn-2t+2, n=3-6, t=2), light straight

chain alkenes (CnH2n, n=3-6), and dienes (CnH2n-2, n=4-6). The reactions on dispersed

H4SiW12O40 clusters (0.075 4 12 40 2

1H SiW O SiO

mmol g ) at 573 K led to constant alkanal conversion

rates and carbon selectivities within experimental errors for reaction times above 155 min, at

141

which stable reactivities were attained, as shown for butanal reactions in Figure 5.1a. Butanal

conversion rates above 155 min remained at 9.1±0.7 mmol (molH+ s)−1 and carbon selectivities

towards C8H14O, C12H20O, C8+ hydrocarbons, C4H8, and C4H6 at 66±1 %, 16.4±0.2 %,

11.2±0.6 %, 0.9±0.05 %, and 4.1±0.3 %, respectively, with the rest (<2 %) being minor (e.g.,

C7H14) or unidentified products. Pyridine adsorption followed by temperature programmed

desorption and infrared spectroscopic studies of pyridine adsorption gave the H+ site densities.

The H+ site densities, expressed as the ratio of H

+ sites to H4SiW12O40 clusters, decreased

drastically by ~40% [from 2.56±0.09 to 1.54±0.32 molH+(molH4SiW12O40)−1] after the initial

exposure to butanal reactant (within 30 min), but remained relatively constant at 1.54±0.32

molH+(molH4SiW12O40)−1 for the longer reaction duration (30-725 min), as shown in Figure 5.1b.

I have confirmed that the polyoxometalate clusters remain stable at 573 K, because treating the

H3PW12O40 in flowing He at 573 K for 3 h only decreases the H+ site density by 6 % (from 3.06

to 2.89 molH+ molcluster−1

, Fig. S5.1, Appendix). Therefore, the marked decrease of the H+ sites

on H4SiW12O40 during butanal reaction was caused predominantly by the formation of coke and

surface acetate (as confirmed from the infrared spectroscopic study reported elsewhere[19]). The

surface acetate was formed through alkanal oxidation by the lattice oxygen on polyoxometalate

clusters, as detected by the appearance of 1580 cm−1

band, which corresponded to the v(OCO)

symmetric vibration detected in the infrared spectra during butanal adsorption on H4SiW12O40 at

348 K [19]. The coke formation was confirmed from the infrared spectra of H4SiW12O40, taken

after butanal reaction at 573 K for 5 min, which showed bands at 1580 cm−1

resulted from the

stretching of the aromatic C=C bond without appearance of the C-H bands (see Fig. S5.2 in

Appendix). The coke formation was also confirmed by temperature programed oxidation (TPO)

carried out on spent H4SiW12O40/SiO2 catalysts, which showed that 5.2 wt.% coke was formed

after butanal reactions for 8 h at 573 K (see Sec. 5.6.3 in Appendix). These coke species were

formed via butanal condensation and sequential cyclization-dehydrogenation reactions on fresh

H4SiW12O40 catalysts.

142

Figure 5.1. (a) Overall butanal (C4H8O) conversion rates (◊) and carbon selectivities for C4H6

(○), C4H8 (∆), C8H14O (▼), C12H20O (■) , and C8+ hydrocarbons (labeled C8+ HC, ●) as a

function of time-on-stream during butanal reactions on H4SiW12O40 clusters at 573 K [butanal

pressure 1.1 kPa, 0.045 +1

butanal Hmol (mol s) , butanal conversion=18-24 %]; (b) H

+ site density,

expressed as the number of H+ site per H4SiW12O40 cluster remained after butanal reactions at

573 K, as a function of time-on-stream [butanal pressure 1.1-4.4 kPa, space velocity 0.045-0.18

+1

butanal Hmol (mol s) ].

Figure 5.2 shows the effects of space velocity on the conversion and carbon selectivities during

butanal deoxygenation on H4SiW12O40 clusters at 623 K. As space velocities decreased from

0.26 to 0.07 +4 8

1C H O H

mol (mol s) , butanal conversions increased from 13.3 % to 23.5 %. As the

conversion increased, the carbon selectivities towards the C8H14O fraction, which contained

above 93 % 2-ethyl-2-hexenal balanced with a small amount of its isomers, decreased from 73.7 %

to 57.6 %, whereas those towards the larger C12H20O, C16H26O, and C8+ hydrocarbons, which

include aromatics (predominantly alkyl benzenes C12H18 and C16H24), cycloalkadienes (C8H12),

and cycloalkenes (C8H14), concomitantly increased (Fig. 5.2a). In contrast, carbon selectivities

towards butene (C4H8) and butadiene (C4H6) remained insignificant (<2.5 %) throughout the

entire range of butanal conversion, as shown in Figure 5.2b.

143

Figure 5.2. Butanal conversions and carbon selectivities to (a) C8H14O (▼), C12H20O (■),

C16H26O (▲), and C8+ hydrocarbons [●, labeled C8+ HC, including C4tH6t aromatics (t=3 or 4),

cycloalkadienes (t=2), and C4tH6t+2 cycloalkenes (t=2)] and (b) C4H6 (○) and C4H8 (∆) during

butanal (C4H8O) reactions on H4SiW12O40 clusters [0.075 4 12 40 2

1H SiW O SiO

mmol g ] as a function

of space velocity at 623 K (1.1 kPa butanal in He, time-on-stream >155 min, at which stable

conversions and selectivities were attained).

These trends of time-dependent rate and selectivity remained the same for other alkanals

(CnH2nO, n=3-6) during their deoxygenation reactions on H4SiW12O40 clusters at 573 K. Figure

5.3 shows the carbon distributions for these different alkanals during their deoxygenation

sojourns at steady-state (time-on-stream=275 min). Across the alkanal homolog, the reactions

produce chemical species with carbon numbers equaled the multiple of that in the parental

alkanal reactants. These selectivity trends indicate that alkanal reactants are incorporated

systematically into the alkanal condensation products in a stepwise growth mechanism, which

consists of a series of consecutive, intermolecular C=C bond formation events. During each of

these C=C bond formation events, an alkanal monomer CnH2nO is incorporated into and a H2O

molecule is ejected from the growing oxygenate molecule formed from multiple alkanal units, as

captured in Scheme 5.1. Large oxygenates (CtnH2tn-2t+2O, n=3-6, t=2-4) are produced; these

oxygenates contain carbon numbers systematically increased by Ctn units (Steps 2a, 3a, and 4a,

144

Scheme 5.1), where subscript n denotes the carbon number of the reactant (n=3-6) and t a

positive integer (t=2-4).

Figure 5.3 (a-d). Carbon distributions of the products, including oxygenates ( , from Steps

2a-2c, 3a-3c, 4a, etc. in Scheme 5.1), aromatics ( , from Steps 3d-3e, etc.), cycloalkadienes

( , from Step 2d), n-dienes ( , from Step 1b), and n-alkenes ( , from Step 1a), during (a)

propanal, (b) butanal, (c) pentanal, and (d) hexanal reactions on H4SiW12O40 clusters at 573 K

[0.075 4 12 40 2

1H SiW O SiO

mmol g , space velocity=0.045 +

1alkanal H

mol (mol s) , alkanal pressure=1.1

kPa, time-on-stream=275 min, conversion=17 %, 30 %, 47 %, and 68 % for propanal, butanal,

pentanal, and hexanal, respectively].

3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 200

5

10

15

75

80

85

C5n

C4n

C3n

C2n

0.8%1%

0.7%0.6%2%

14%

1%

80%

Carb

on d

istr

ibution (

%)

Carbon number

Oxygenates

Aromatics

Cyclo-dienes

n-Dienes

n-Alkenes

0.4%

Cn

3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 200

5

10

15

75

80

85

C5n

C4n

C3n

C2n

0.8%1%

0.7%0.6%2%

14%

1%

80%

Carb

on d

istr

ibution (

%)

Carbon number

Oxygenates

Aromatics

Cyclo-dienes

n-Dienes

n-Alkenes

0.4%

Cn

3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 200

5

10

15

75

80

85

C5n

C4n

C3n

C2n

0.8%1%

0.7%0.6%2%

14%

1%

80%

Carb

on d

istr

ibution (

%)

Carbon number

Oxygenates

Aromatics

Cyclo-dienes

n-Dienes

n-Alkenes

0.4%

Cn

3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 200

5

10

15

75

80

85

C5n

C4n

C3n

C2n

0.8%1%

0.7%0.6%2%

14%

1%

80%

Ca

rbo

n d

istr

ibutio

n (

%)

Carbon number

Oxygenates

Aromatics

Cyclo-dienes

n-Dienes

n-Alkenes

0.4%

Cn

3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 200

5

10

15

75

80

85

C5n

C4n

C3n

C2n

0.8%1%

0.7%0.6%2%

14%

1%

80%

Carb

on d

istr

ibution (

%)

Carbon number

Oxygenates

Aromatics

Cyclo-dienes

n-Dienes

n-Alkenes

0.4%

Cn

145

Scheme 5.1 also captures the sequential reactions of the alkenals (CtnH2tn-2t+2O, n=3-6, t=2-3).

These alkenals may undergo sequential ring closure, evolving cycloalkenols or cycloalkadienols

(CtnH2tn−2t+2O, Steps 2b and 3b). The cycloalkenols/cycloalkadienols could either isomerize to

cycloalkanones or cycloalkenones (CtnH2tn−2t+2O, Steps 2c and 3c) or dehydrate to form

cyclodienes or aromatics (CtnH2tn−2t, Steps 2d and 3d). The numbers of carbon atoms in the

cyclodiene or aromatic product fractions are multiples of the alkanal monomers (tn) and related

to the number of chain growth events, which equals t−1 (where t=2-4). Transalkylation of these

Ctn hydrocarbons that forms hydrocarbons with tn−1 or tn+1 carbon atoms (Ctn−1 and Ctn+1) may

occur, as reported for alkanal reactions on zeolites [8, 11, 18], but the reaction is kinetically

insignificant on polyoxometalate clusters, as confirmed from the low carbon selectivities of <1%

on H4SiW12O40 clusters (to be discussed in Section 5.3.4). The primary aromatic products (e.g.,

alkyl benzenes) can undergo dehydrogenation (Step 3e) and evolve alkyl tetralin species (e.g.,

C3nH6n−8, n=4-6), which are active hydrogen donors required for the light alkene (CnH2n)

formation, as discussed later in this section and Section 5.3.2. I rule out both decarbonylation and

decarboxylation as the predominant routes for oxygen removal, because COx formation was

below the detectable amounts (<0.03% carbon selectivity) over a broad range of temperature

(473-673 K) and residence time [0.045-0.27 +1

alkanal Hmol (mol s) ].

146

Scheme 5.1. Pathways for alkanal (CnH2nO) chain growth resulting in larger alkenals (and their

isomers, CtnH2tn-2t+2O, n=3-6, t=2-3) and hydrocarbons (including cycloalkadienes and aromatics,

CtnH2tn-2t, n=3-6, t=2-3) (R, R1, and R2 represent either alkyl group or H).

In parallel to the condensation and chain growth reactions, alkanals (CnH2nO, n=3-6) also convert

to light alkenes (CnH2n, n=3-6) and dienes (CnH2n-2, n=4-6) (Steps 1a and 1b, respectively,

Scheme 5.1). Over 90 % of the light alkenes and dienes in their respective product fractions (Fig.

5.3) are molecules with carbon numbers identical to their parental alkanal reactants (CnH2n/

2 1

22C H

m n

m mm

>90 % and CnH2n-2/2 1

2 22C H

m n

m mm

=100 % from CnH2nO alkanal reactions,

n=3-6, Fig. 5.3). Such near exclusive selectivities towards these specific alkenes and dienes

provide the evidence of their direct formation in reactions that must involve direct hydrogen

transfer-dehydration and isomerization-dehydration, respectively, without chain lengthening or

contraction.

The space velocity effects (Fig. 5.2) and product distributions (Fig. 5.3) led to the proposed

reaction network in Scheme 5.2, which appears to be general for the deoxygenation of linear

alkanals (CnH2nO, n=3-6) on H4SiW12O40 clusters. Two alkanal molecules undergo a primary,

intermolecular C=C bond formation step via an aldol condensation-dehydration route (e.g.,

butanal condensation forms 2-ethyl-2-hexenal, Cycle 1) that results in a dimeric alkenal (CtnH2tn-

147

2t+2O, n=3-6, t=2, labeled “Dimer”) [11]. The alkenal (CtnH2tn-2t+2O) reacts with another alkanal

reactant (Cycle 1.1) in a sequential intermolecular C=C bond formation step and, as a result,

augments its molecular size, thus evolving a larger alkenal (CtnH2tn-2t+2O, n=3-6, t=3, labeled

“Trimer”). The intermolecular C=C bond formation step may occur again (Cycle 1.2), leading to

an even larger alkenal (CtnH2tn-2t+2O, n=3-6, t=4, labeled “Tetramer”). These sequential chain

growth events are captured by Equations 5.1a to 5.1c below:

CnH2nO+ CnH2nO C2nH4n-2O+H2O (5.1a)

C2nH4n-2O + CnH2nO C3nH6n-4O+H2O (5.1b)

C3nH6n-4O + CnH2nO C4nH8n-6O+H2O (5.1c)

and more generally:

tCnH2nO CtnH2tn-2t+2O +(t−1)H2O (5.1d)

as also depicted in Scheme 5.1, Steps 2a, 3a, and 4a, respectively. The dimeric oxygenates

(C2nH4n-2O, n=3-6) may undergo concurrent reactions and form cycloalkadienes (C2nH4n-4, n=3-6,

t=2) via cyclization-dehydration and isomerization reactions (Step Cyclization 1). Similarly,

larger oxygenates derived from trimeric and tetrameric species (CtnH2tn−2t+2O, n=3-6, t=3-4) may

undergo cyclization-dehydration (Steps Cyclization 2 and Cyclization 3) or cyclization-

dehydration-dehydrogenation (Step Cyclization 2.1) reactions that strip all of their oxygen atoms

and evolve as aromatic species [including substituted benzenes (CtnH2tn−2t) and tetralins

(CtnH2tn−2t−2), n=3-6, t=3-4].

148

Scheme 5.2. Reaction network for alkanal (CnH2nO, n=3-6) deoxygenation on H4SiW12O40

clusters capturing the intermolecular C=C bond formation (Cycle 1, 1.1, and 1.2), intramolecular

C=C bond formation (Cycle 2), isomerization-dehydration (Cycle 3 and 3.1), the secondary

cyclization-dehydration and dehydrogenation reactions (Cyclization 1, 2, 2.1, and 3), illustrated

with butanal as an example.

The primary, intermolecular C=C bond formation event (Cycle 1) is kinetically coupled with

separate, interdependent catalytic cycles, also shown in Scheme 5.2, of: (i) direct deoxygenation

of alkanal (Cycle 2, Intramolecular C=C formation) that evolves light alkene, during which an

intramolecular C=C bond is created, first by H transfer from a H-donating agent to a protonated

alkanal, followed by dehydration and (ii) competitive isomerization-dehydration via a

bimolecular pathway (Cycle 3, Isomerization-dehydration) that forms light dienes (CnH2n-2, n=4-

6). These reactions are:

149

CnH2nO+R’H2 CnH2n+H2O+R’ (R’H2 represents a H-donor) (5.2a)

CnH2nOCnH2n-2+H2O (5.2b)

The alkanal isomerization-dehydration could also proceed via a monomolecular route (Cycle 3.1,

Isomerization-dehydration, Scheme 5.2), which is not kinetically coupled to the pathway of

intermolecular C=C bond formation.

Similar deoxygenation pathways are found for C3-C6 alkanals (n=3-6), but their relative rates and

selectivities vary systematically with the carbon number n, because of the difference in

molecular dimensions, heats of adsorption, and hydride ion affinities, through their influences on

the kinetic properties (i.e., activation free energies) of these different pathways. The evidence of

catalytic couplings and the mechanistic details of the primary reactions are provided next in

Sections 5.3.2 and 5.3.3.

5.3.2. Kinetic coupling of alkanal chain growth, deoxygenation, and isomerization-dehydration cycles

I describe next the kinetic couplings of alkanal chain growth (Cycle 1), intramolecular C=C bond

formation (Cycle 2), and isomerization-dehydration (Cycle 3) catalytic cycles shown in Scheme

5.2. The reaction stoichiometry for intramolecular C=C bond formation dictates that each alkene

formation requires two H atoms (CnH2nO+2HCnH2n+H2O). These hydrogen atoms, in the

absence of external H sources, must come from either the reaction intermediates or products. A

carbon and hydrogen balance carried out in our previous studies on propanal reactions catalyzed

by H-MFI zeolite[11] identifies that aromatic products or precursors to aromatics, typically

substituted cyclohexadienes and tetralins produced from the secondary cyclization reactions of

larger, unsaturated alkenals (Steps Cyclization 1, Cyclization 2, Cyclization 2.1, and Cyclization

3, e.g., 5,6-dimethyl-1,3-cyclohexadiene and 5,7-dimethyl-tetralin produced in butanal reactions),

are the H donors (R’H2) in Equation 5.2a above. These substituted cycloalkadienes or tetralins

increase their extent of unsaturation via H donation, forming more stable alkyl benzenes or

naphthalenes with delocalized π-bond [21]. The catalytic involvement of substituted

cycloalkadienes and tetralins was evident from an increase in butene formation rates with

cyclohexadiene and tetralin incorporation during butanal reactions on H-FAU zeolites (at 573 K),

as reported elsewhere [18]. Indeed, turnover rates for intramolecular C=C bond formation ( Intrar )

150

increased linearly with the pressure of the cyclic hydrocarbon products (including

cycloalkadienes and aromatics), HCP , for all C3-C6 alkanals, as shown in Figure 5.4a and

captured by Equation 5.3:

Intra HCIntra,effr k P (5.3)

where 𝑘Intra,eff is the effective rate constant for intramolecular C=C bond formation. The H

balance[11] and the direct correlations between the rates 𝑟Intra and cyclic hydrocarbon pressures

(Fig. 5.4a) led us to propose that the inter- and intramolecular C=C bond formation cycles

(Cycles 1 and 2, Scheme 5.2) are kinetically coupled via the intermolecular hydrogen transfer

step.

Figure 5.4. (a) Rates for intramolecular C=C bond formation ( Intrar , Cycle 2 in Scheme 5.2) as a

function of the total pressure of C8-C16 cyclic hydrocarbon product fraction ( HCP , including

cycloalkadienes and aromatics), (b) rates for intermolecular C=C bond formation ( Interr , Cycle 1

in Scheme 5.2) as a function of alkanal pressure ( alkanalP , average alkanal pressure), and (c) rates

for isomerization-dehydration ( Dehyr , Cycle 3 and Cycle 3.1 in Scheme 5.2) as a function of

alkanal pressure ( alkanalP , average alkanal pressure) during the reactions of alkanals [CnH2nO,

n=3-6; propanal (▲), butanal (■), pentanal ( ), and hexanal ( )] on H4SiW12O40 clusters [573

K, 0.045-0.44 +1

alkanal Hmol (mol s) , time-on-stream=275-600 min, alkanal conversion=14-17 %,

26-31 %, 45-47%, and 68-72 % for propanal, butanal, pentanal, and hexanal, respectively]

151

Next, I probe the catalytic coupling of the initial intermolecular C=C bond formation (Cycle 1)

and isomerization-dehydration (Cycle 3) reactions shown in Scheme 5.2, which form 2-ethyl-2-

hexenal and butadiene, respectively. The rates for intermolecular C=C bond formation ( Interr )

and isomerization-dehydration ( Dehyr , of both Cycle 3 and Cycle 3.1) measured for butanal

reactions on H4SiW12O40 clusters at 573 K are shown in Figure 5.5 as a function of time-on-

stream. The rates Interr increase while those of Dehyr concomitantly decrease during the first 155

min of reaction before approaching constant values [ Interr =3.8±0.2 +1

Hmmol (mol s) and Dehyr

=0.40±0.07 +1

Hmmol (mol s) ]. Despite changes in these individual rate values during the initial

times (<155 min), their combined rates ( Dehyr + Interr ) remain relatively constant [4.4±0.5

+1

Hmmol (mol s) ], an indication that the catalytic cycle of intermolecular C=C bond formation

(Cycle 1) is kinetically coupled with at least one of the isomerization-dehydration pathways

(Cycle 3 and/or Cycle 3.1), possibly through sharing common surface intermediates. It has been

proposed that alkanal dehydration that forms alkadiene proceeds via an allylic alcohol

intermediate upon alkanal isomerization [17]. The alkanal isomerization involves an initial

protonation of the carbonyl group at a H+ site and a sequential β-hydrogen abstraction by a

vicinal base site, as proposed for 2-methyl-butanal isomerization on BPO4 and AlPO4 catalysts

[17]. These previous studies led us to further postulate that acid-base site pairs are required for

alkanal isomerization-dehydration (Cycle 3 and Cycle 3.1). The anion of tungstosilicic acid

cluster (SiW12O404−

) has multiple oxygen sites (both bridging and terminal oxygen atoms) that

can act as the base sites [22]. As butanal reactions proceed, larger carbonaceous species (e.g.,

cokes) were deposited on the H4SiW12O40 clusters, covering the oxygen sites. In contrast, the H+

site density remained relatively stable for reaction times above 30 min (Fig. 5.1b), probably

because the alkanals adsorbed at the H+ sites and protected them from being covered by cokes.

The decrease in base sites prevents these shared intermediates, the precursors for the

isomerization-dehydration and intermolecular C=C bond formation, from undergoing

isomerization-dehydration (Cycle 3 and Cycle 3.1) and in turn promotes the competitive

intermolecular C=C bond formation (Cycle 1), as shown in Figure 5.5, to be discussed in details

in Section 5.3.3 and 5.3.4.

152

Figure 5.5. Rates for intermolecular C=C bond formation ( Interr , ■, Cycle 1 in Scheme 5.2) and

isomerization-dehydration ( Dehyr , ▲, Cycle 3 and Cycle 3.1 in Scheme 5.2) and the combined

rate ( Dehyr + Interr , ○) as a function of time-on-stream on H4SiW12O40 catalysts (0.075

4 12 40 2

1H SiW O SiO

mmol g ) at 573 K [space velocity=0.063 +

1butanal H

mol (mol s) , butanal

conversion=18-24 %].

5.3.3. Mechanisms for the formation of alkenals, alkenes, and dienes via primary alkanal reactions on H4SiW12O40

clusters

Scheme 5.3 shows a proposed sequence of elementary steps for the primary alkanal reactions,

which includes intermolecular C=C bond formation (Steps G1, A1, and R1.1-R1.5),

intramolecular C=C bond formation (Steps A1 and R2.1-R2.4), and isomerization-dehydration

via both the bimolecular (Steps A1, R1.1, and R3.1a-R3.2a) and monomolecular (Steps A1 and

R3.1b-R3.2b) pathways. Pseudo steady-state treatments of all reactive intermediates in this

sequence and the assumption of protonated butanals as the most abundant surface intermediates

lead to a set of rate equations consistent with the observed kinetic dependencies for the primary

reactions shown in Figure 5.4 and with the coupling of the various catalytic cycles in Scheme 5.2,

as described next.

153

Scheme 5.3. Mechanism for intermolecular C=C bond formation (Steps G1, A1, and R1.1-R1.5,

also shown as Cycle 1 in Scheme 5.2), intramolecular C=C bond formation (Steps A1 and R2.1-

R2.4, also shown as Cycle 2 in Scheme 5.2), and isomerization-dehydration via bimolecular

pathway (Steps R3.1a-R3.2a, also shown as Cycle 3 in Scheme 5.2) and monomolecular pathway

(Steps R3.1b-R3.2b, also shown as Cycle 3.1 in Scheme 5.2) during alkanal reactions on

H4SiW12O40 clusters (R=H, CH3, C2H5, and C3H7 for propanal, butanal, pentanal, and hexanal,

respectively; R’H2 represents a H-donor).

Alkanals adsorb on H+ sites as protonated alkanals (CnH2nOH

+, Step A1, equilibrium constant

Kads) [23], their adsorbed alkenol tautomers (Step A3.1) [24-26], or in their physisorbed

154

equivalence via hydrogen bonding between their oxygen atoms to the H+ sites (Steps A2.1-A2.2)

[23, 24, 27]. The alkanals and their various surface species are likely in chemical equilibrium, as

inferred from previous infrared studies of acetone adsorption on H-ZSM-5[23] that detected both

the carbonyl group (1658-1671 cm−1

) and the alcohol species (1375 cm−1

and 880 cm−1

for in-

plane and out-of-plane O-H bending, respectively). The adsorption weakens the O-H bridging

bands of the zeolite at 2850 cm−1

and 2380 cm−1

to different extents as a result of their

interactions with the physisorbed and protonated carbonyl groups, respectively.

Protonated carbonyl groups were evident in infrared spectroscopic studies from a red shift of the

v(C=O) band in acetone during its adsorption from 1720 cm-1

on Na-ZSM-5 to 1671-1658 cm−1

on H-ZSM-5 [23]. Adsorbed enols, the surface tautomers, were confirmed from H-D exchange

between the H in the adsorbed 13

C-2-acetone and the D+ site and between the D in the adsorbed

acetone-d6 and the H+ site in ZSM-5 zeolites at ambient temperature [25]. They were also

evident from the appearance of the signal at ~180 ppm in 13

C NMR spectra as acetone was

adsorbed on H-ZSM-5 and H-Y zeolites between 298 K and 453 K [26]. DFT calculations gave

an activation barrier of 75.3 kJ mol−1

and reaction energy of 31.0 kJ mol−1

for the keto-enol

tautomerization of acetaldehyde on the H+ sites in H-ZSM-5 zeolites [24], thus protonated enol is

much less stable than the protonated alkanal and is a less abundant surface intermediate. It is

noted that the keto-enol tautomerization on H+ sites (Step A3.1) has a high activation barrier (e.g.,

∆‡Henol= 75.3 kJ mol

−1 for acetaldehyde tautomerization on H-ZSM-5);[24] the barrier is even

higher in the gas phase (e.g., ∆‡Henol= 285 kJ mol

−1 for acetone tautomerization in the gas phase)

[28]. The activation enthalpy for keto-enol tautomerization (Step A3.1, ∆‡Henol) may be much

higher than that for aldol condensation (Step R1.1-R1.3, ∆‡Haldol). For example, the activation

enthalpies for aldol condensation are 8.4 kJ mol−1

for 2-propenol and formaldehyde reactions on

H-ZSM-5[28] and 27 kJ mol−1

for 1-butenol and butanal reactions on Ti-OH sites [29], based on

DFT calculations. The bimolecular aldol condensation step (Steps R1.1-R1.3), however, has a

more negative activation entropy than the monomolecular keto-enol tautomerization step (Step

A3.1). In addition, the aldol condensation (Steps R1.1-R1.3) involves the gaseous enol tautomers

with low concentration, whereas the keto-enol tautomerization (Step A3.1) occurs on the

protonated alkanals, the most abundant surface species. As a result, Steps R1.1-R1.3 have a

lower rate than Step A3.1. For these reasons, the aldol condensation step remains as the rate

limiting step for the intermolecular C=C bond formation, and exhibits first order rate dependence

155

to alkanal pressure, as shown later in this section. Since the alkenol formation via the reversible

keto-enol tautomerization (Steps A3.1) is not a rate-limiting step, and the alkanal adsorption

(Step A1) and alkenol desorption (Step A3.2) steps are all reversible, the reversible keto-enol

tautomerization in the gas phase (Step G1) would reach chemical equilibrium during steady-state

reactions.

A fraction of the adsorbed alkanals on strong H+ sites underwent oxidation on the H4SiW12O40

clusters and formed stable surface acetate (not shown in Scheme 5.3). These acetate species were

evident from the appearance of the v(OCO) symmetric vibrational stretching band at 1580 cm−1

during exposure of H4SiW12O40 clusters to butanal at 348 K; the density of surface acetate

increased proportionally with the number of strong H+ sites contained within the clusters, as

determined by pyridine-TPD [19], an indication that these acetate species were formed only on

strong H+ sites. During alkanal reaction at 573 K, these strong H

+ sites were poisoned by the

formation of coke (Fig. S5.1, Appendix), as confirmed from the disappearance of the pyridine

desorption peak at 720 K during pyridine-TPD of the spent H4SiW12O40 catalysts (see Fig. S5.5,

Appendix). Therefore, surface acetate species were not present during steady-state reactions at

573 K. The chemical titration of the H+ sites with C3-C6 alkanals at 348 K, a much lower

temperature than that of the steady-state reaction, shows near unity alkanal-to-H+ site ratios

(CnH2nO:H+= 1.15, 1.0, 1.08, and 1.05 for n=3, 4, 5, and 6, respectively), an indication that the

catalytically active H+ sites not poisoned by cokes are occupied predominantly by protonated

alkanal monomers (Step A1, labeled A* in Scheme 5.3) and their isomers (protonated alkenols

and physisorbed alkanals), which are in equilibrium with each other (Steps A3.1 and A2.2), as

shown in the shaded area in Scheme 5.3.

The intermolecular C=C bond formation cycle (Cycle 1 of Scheme 5.2) is initiated by keto-enol

tautomerization, which transforms a small amount of alkanals to alkenols (Step G1, Scheme 5.3).

As depicted in Scheme 5.3, I postulate that the alkenol can co-adsorb on the lattice oxygen site

adjacent to a protonated alkanal and form a co-adsorbed alkanal-alkenol pair (labeled AAP*,

Step R1.1). Within the AAP* intermediate, nucleophilic attack of the C=C bond in the co-

adsorbed alkenol (CnH2n-1OH*, labeled Alkenol*) to the adjacent carbonyl carbon of the

protonated alkanal (CnH2nOH+) (Step R1.2) via a bimolecular transition state [labeled TS(C-C)*]

leads to the formation of an aldol C2nH4nO2 (Steps R1.3-R1.4, labeled “Aldol”), which upon

sequential dehydration (Step R1.5) forms the alkenal C2nH4n-2O. The rate equation for

156

intermolecular C=C formation, derived by considering the nucleophilic attack (Step R1.2) as a

kinetically-relevant step, is:

AAP ads alkanal taut alkanal

C-C

-AAP C-C iso,bi

Inter


ads alkanal

-AAP C-C iso,bi

+

+

+ +

+

[H ]

[H ][H ] [H ]

(H site) (A*)

k K P K Pk

k k kr

k K P K PK P

k k k

(AAP*) (other)

(5.4)

where adsK and tautK are the equilibrium constants for alkanal adsorption (Step A1) and keto-

enol tautomerization (Step G1), respectively; terms AAPk and -AAPk are the rate constants for the

forward and reverse steps of AAP* formation (Step R1.1), respectively; C-Ck is the rate constant

for aldol condensation (Step R1.2) and iso,bik for the bimolecular pathway of alkanal

isomerization (Step R3.1a). [H+] denotes the H

+ site concentration whereas alkanalP denotes the

alkanal pressure. These rate and equilibrium constants are defined in Scheme 5.3. The relative

magnitudes of the denominator terms ads alkanal

+[H ]K P ,

AAP ads alkanal taut alkanal -AAP C-C iso,bi

1+[H ]k K P K P k k k

, and reflect the relative coverages of

protonated alkanals (A*), the alkanal-alkenol pairs (AAP*), and other adsorbed species (e.g.,

alkanal isomers and alkenal products), respectively. The term ads alkanal

+[H ]K P is the predominant

term in the denominator, when protonated alkanals (A*) are the most abundant surface

intermediates, as confirmed from the chemical titration with C3 to C6 alkanals. Therefore,

Equation 5.4 simplifies to rates that vary linearly with alkanal pressure:

AAP C-C tautInter alkanal Inter,eff alkanal

-AAP C-C iso,bi

k k Kr P k P

k k k

(5.5a)

AAP C-C tautInter,eff

-AAP C-C iso,bi

k k Kk

k k k

(5.5b)

This first-order dependence for intermolecular C=C bond formation on alkanal pressures

(𝑃alkanal) is confirmed for alkanal (CnH2nO, n=3-6) reactions on H4SiW12O40 clusters at 573 K

157

during steady state, as shown in Figure 5.4b. The term AAP C-C taut -AAP C-C iso,bi

1k k K k k k

in

Equation 5.5b reflects the effective rate constant, Inter,effk . I determine the rate constants (to be

discussed in Sec. 5.3.5) from the rate data using an integral reactor model, because alkanal

conversions were higher than 10 % (14-17 %, 26-31 %, 45-47 %, and 68-72 % for propanal,

butanal, pentanal, and hexanal, respectively).

The protonated alkanal may undergo an alternative catalytic cycle (Cycle 2, Scheme 5.2), during

which it accepts a hydride ion from a hydride donor (denoted R’H2) in a kinetically-relevant step

(Step R2.1) to form a carbonium ion transition state with a shared hydride ion, followed by its

decomposition (Step R2.2) and kinetically irrelevant dehydration (Step R2.3) to evolve the n-

alkene, as captured in Scheme 5.3. Proton donation from the carbenium ion of the hydride donor

(R’H+) to the polyoxometalate surfaces (Step R2.4) completes the catalytic cycle and regenerates

the H+ site. This hydride transfer mechanism is similar to those previously studied with DFT

calculations between alkanes and protonated alkenes on H3Si-OH-AlH2-O-SiH3 clusters[30-32]

and between alkanes to alkoxides in mordenite zeolites [33]. This mechanism leads to the

following rate equation for intramolecular C=C bond formation ( Intrar ):

2HT ads alkanal R'H

Intra


ads alkanal

-AAP C-C iso,bi

+

+

+ +

+

[H ]

[H ][H ] [H ]

(H site) (A*) (AAP*) (

k K P Pr

k K P K PK P

k k k

other)

(5.6)

The denominator terms in this rate equation remain identical to those in Equation 5.4, because

both the inter- and intramolecular C=C bond formation reactions occur at the same sites. HTk

represents the rate constant for the hydride transfer step (Step R2.1) and 2R'HP is the partial

pressure of the hydride donors. When protonated alkanal (A*) is the most abundant surface

intermediate, Equation 5.6 simplifies to rates that depend strictly on the hydride donor pressure:

2Intra HT R'Hr k P (5.7)

In the absence of external hydrogen sources, the large cyclic hydrocarbon products including

aromatics (e.g., alkyl benzenes and alkyl tetralins) and precursors to aromatics (e.g.,

158

cycloalkadienes) are the only hydride donors, as they undergo dehydrogenation to further

increase their extent of unsaturation. The reactivities of these cyclic hydrocarbon products as H-

donors vary with their chemical identities, as demonstrated in our previous work [18]. Assuming

that x is the fraction of a specific H-donor R’H2(x) within the cyclic hydrocarbon products

(where subscript x denotes the chemical identity) and 2HT,R'H ( )xk is the rate constant for transfer

hydrogenation between R’H2(x) and the protonated alkanal, the rate for intramolecular C=C bond

formation, Intrar , is:

2 2 2Intra HT, HT, HC HCR'H ( ) R'H ( ) R'H ( ) Intra,eff1 1 xx x x

t t

x xr k P k P k P

(5.8a)

2HT,Intra,eff R'H ( )1 xx

t

xk k

(5.8b)

According to Equation 5.8a, the rate of intramolecular C=C bond formation ( Intrar ) increases

linearly with the partial pressures of cyclic hydrocarbon products ( HCP ), consistent with the rate

dependences shown in Figure 5.4a. Intra,effk is the effective rate constant for intramolecular C=C

bond formation and its value depends on the fractions ( x ) and the hydride transfer rate

constants (2HT,R'H ( )xk ) of the various H-donors [R’H2(x), x=1,2,…], as shown in Equation 5.8b.

This direct, linear relation between the rates of intramolecular C=C bond formation and the

hydrocarbon pressures is found not only for C3-C6 alkanal reactions on H4SiW12O40 clusters at

573 K (Fig. 5.4a), but also for butanal reactions on H-MFI and H-FAU zeolites [18].

Another parallel catalytic cycle (Cycle 3, Scheme 5.2) converts the alkanal to n-diene via the

decomposition of an adsorbed alkanal-alkenol pair (AAP*) into a protonated alkanal (A*) and an

allylic alcohol (Step R3.1a), followed by the dehydration of the allylic alcohol (Step R3.2a) to

evolve a diene, as shown in Scheme 5.3. The relative rates for isomerization and dehydration are

probed by comparing the rates of butadiene formation from butanal ( Dehy,butanalr ) and those of

crotyl alcohol dehydration ( Dehy,crotyl alcoholr ) on H4SiW12O40 clusters at 573 K. The butadiene

formation rates are significantly smaller (by 118 times) than those of crotyl alcohol dehydration

[ Dehy,butanalr = 0.38±0.06 +1

Hmmol (mol s) vs. Dehy,crotyl alcoholr = 45 +

1

Hmmol (mol s) , 573 K,

159

space velocity of 0.045 +1

reactant Hmol (mol s) ], thus the initial isomerization step (Step R3.1a)

must be kinetically relevant. It is expected that this kinetically relevant allylic alcohol formation

step could also take place on a protonated alkanal monomer (A*) (Steps R3.1b). The proposed

mechanism leads to the rate equations for diene formation via bimolecular (Cycle 3, Dehy,bir ) and

monomolecular (Cycle 3.1, Dehy,monor ) pathways, respectively:


iso,bi

-AAP C-C iso,bi

Dehy,bi


ads alkanal

-AAP C-C iso,bi

+

+

+ +

+

[H ]

[H ][H ] [H ]

(H site) (A*)

k K P K Pk

k k kr

k K P K PK P

k k k

(AAP*) (other)

(5.9)

iso,mono ads alkanal

Dehy,mono


ads alkanal

-AAP C-C iso,bi

+

+

+ +

+

[H ]

[H ][H ] [H ]

(H site) (A*) (AAP*)

k K Pr

k K P K PK P

k k k

(other)

(5.10)

where iso,bik and iso,monok are the rate constants for alkanal isomerization via bimolecular (Step

R3.1a) and monomolecular (Step R3.1b) pathways, respectively. All the other rate and

equilibrium constants are defined above and in Scheme 5.3 and the denominator terms remain

identical to those in Equations 5.4 and 5.6. Equations 5.9 and 5.10 simplify to the followings,

when protonated alkanal (A*) is the most abundant surface intermediate:

tautAAPiso,bi

Dehy,bi alkanal Dehy,bi,eff alkanal

-AAP C-C iso,bi

k k Kr P k P

k k k

(5.11a)

tautAAPiso,bi

Dehy,bi,eff

-AAP C-C iso,bi

k k Kk

k k k

(5.11b)

Dehy,mono iso,mono Dehy,mono,eff alkanal0r k k P (5.12a)

Dehy,mono,eff iso,monok k (5.12b)

160

Equations 5.11a and 5.12a show that 𝑟Dehy,bi increases proportionally with alkanal pressure

( alkanalP ) whereas Dehy,monor remains insensitive to alkanal pressure. The terms

tautAAP -AAP C-Ciso,bi iso,bi

1k k K k k k

and iso,monok are the effective rate constants for

bimolecular and monomolecular pathways for alkanal dehydration ( Dehy,bi,effk and Dehy,mono,effk ),

respectively, as shown in Equations 5.11b and 5.12b. The sum of these rates from the two

pathways gives the overall rate of diene formation, Dehyr :

Dehy Dehy,bi Dehy,mono Dehy,bi,eff alkanal Dehy,mono,effr r r k P k (5.13)

In the proposed mechanism (Scheme 5.3), alkanal isomerization (Steps R3.1a and R3.1b)

requires participation of a vicinal lattice oxygen (basic site) adjacent to the H+ site as a hydrogen

abstractor to promote the double bond shift from C=O bond to C=C bond [17]. The rate of

isomerization is expected to depend on the number of these H+ and lattice oxygen pairs that

function together as the acid-base site pairs. The catalytic involvement of an acid-base site pair in

alkanal isomerization was also found on H-ZSM-5 zeolite: the removal of extra-framework

Al2O3 (by ~65 %) from H-ZSM-5 via (NH4)2SiF6 treatment suppresses the rate of isomerization-

dehydration commensurately by about half without affecting those of inter- and intramolecular

C=C bond formation (see Table S5.2 and Fig. S5.4 in Sec. 5.6.4, Appendix). This confirms that a

portion of the isomerization-dehydration turnovers requires extra-framework Al2O3, which

contains a bridging oxygen ion between two Al ions [34], acting cooperatively with an adjacent

H+ site as an acid-base site pair in the isomerization step. On the Keggin anion (SiW12O40

4−), all

the lattice oxygen atoms on the external surface, which include 12 terminal oxygens (W=O) and

24 bridging oxygens (W-O-W), can act as the conjugated base sites for the protons (H+) [22].

Because each terminal oxygen has four and each bridging oxygen has six adjacent surface

oxygen sites, alkanal molecules can easily find the acid-base site pairs to undergo isomerization-

dehydration on H4SiW12O40 clusters. In the contrasting case of MFI zeolites, only the O atoms of

the framework AlO4 tetrahedrons can act as the conjugated base sites. Typically three out of the

four O atoms in AlO4 tetrahedrons are exposed in the zeolite channels and remained accessible to

the reactant. A small amount of extra-framework Al-O-Al structures may be located next to the

H+ sites, in which case the extra-framework O and the H

+ sites form an acid-base site-pair.

161

Generally, the H+ sites in zeolites have much less adjacent base sites than those on H4SiW12O40

clusters. As a result, alkanal molecules have a smaller chance to encounter acid-base site pairs

and thus a lower rate for isomerization-dehydration on zeolites than on H4SiW12O40 [e.g., Dehyr =

1.3 vs. 0.19 +1

Hmmol (mol s) on H4SiW12O40 and H-MFI, respectively, time-on-stream=35 min,

573 K, 1.1 kPa butanal, as shown in Fig. 5.5 and Fig. S5.2c, respectively].

As the reaction time increased, the rate for alkanal dehydration ( Dehyr ) gradually decreased: from

1.3 +1

Hmmol (mol s) at 35 min to 0.44 +

1

Hmmol (mol s) at 155 min (Fig. 5.5), although the

number of H+ sites remained relatively stable during this period (>30 min, Fig. 5.1b). I surmise

that as the reaction time increases, the deposition of heavy products (e.g., aromatics and cokes)

may poison the lattice oxygen (basic sites), making a large portion of the acid-base site pairs

inaccessible to alkanals, thus leading to decrease in iso,bik and iso,monok , and in turn the

isomerization rates (Eqns. 5.11a and 5.12a).

As iso,bik decreases, the rate of AAP* consumption by Step R3.1a decreases and, as a result, the

coverage of AAP* intermediates (Scheme 5.3) increases concomitantly. This trend between

iso,bik and the coverage of the AAP* intermediate is captured by the term

tautAAP -AAP C-Cads alkanal alkanal iso,bi

1+[H ]k K P K P k k k

in the denominator of Equation 5.4.

Such changes led to an increase in the effective rate constant Inter,effk and the rate Interr for

intermolecular C=C bond formation (Cycle 1, Scheme 5.2) with reaction time, according to

Equations 5.5b and 5.5a, as shown in Figure 5.5. After 155 min of reaction, the changes in Interr

and Dehyr values became much slower than during the initial times: the changes in rate value with

time, Interr (∆time-on-stream)−1

and Dehyr (∆time-on-stream)−1

, were one order of magnitude

smaller at 155 min than at 35 min (Fig. 5.5). At these longer reaction times, most of the basic

sites were deactivated by coke deposition and the number of acid and basic sites remained

essentially constants. Thus, iso,bik , iso,monok and the effective rate constants ( Dehy,bi,effk and

Dehy,mono,effk ) remained constant. The rate data in Figure 5.4c reflect the overall rates of diene

162

formation, 𝑟Dehy, at these longer reaction times. The slopes in Figure 5.4c are related to the

effective rate constants for the bimolecular pathway ( Dehy,bi,effk ) while the intercepts reflect the

effective rate constants for the monomolecular pathway ( Dehy,mono,effk ) for diene formation.

5.3.4. Catalytic sequences for secondary cyclization-dehydration that form aromatics and cycloalkadienes on H4SiW12O40

clusters

In this section, I describe the secondary cyclization-dehydration reactions (Cyclization 1,

Cyclization 2, Cyclization 2.1, and Cyclization 3 in Scheme 5.2) that form cycloalkadienes

(CtnH2tn-2t, n=3-6, t=2) and aromatics [CtnH2tn-2t, n=3-6, t=3-4, including alkyl benzenes (t=3) and

alkyl tetralins (t=4)], as shown in Scheme 5.2. These secondary, larger products all contain

carbon numbers equal to multiples (t) of the parental alkanal unit (Cn), as shown from the carbon

distributions in Figure 5.3, because they are products of the stepwise alkanal addition reactions.

In each of these steps, an alkanal unit was added to the carbon chain in an intermolecular carbon-

carbon bond formation event. The carbon chain growth is consistent with the increase in the

carbon selectivities towards larger alkenals (C12H20O and C16H26O) and hydrocarbons (including

C8 cycloalkadienes, C12 and C16 aromatics) as the space velocity decreases during butanal

(C4H8O) reactions at 623 K (Fig. 5.2a). As shown in Scheme 5.1, the larger Ctn alkenal products

(CtnH2tn-2t+2O, n=3-6, t=2-4) may further undergo cyclization to form Ctn cycloalkenols or

cycloalkadienols (CtnH2tn-2t+2O, n=3-6, t=2-4) (Steps 2b and 3b), before their dehydration to

produce Ctn cycloalkadienes (CtnH2tn-2t, n=3-6, t=2) or aromatics (CtnH2tn-2t, n=3-6, t=3-4) (Steps

2d and 3d), as also captured in Equation 5.14:

cyclization dehydration2 2 2 2 2 2 2 2 2C H O C H O C H H O

(alkenal) (cycloalkenol or (aromatic or

tn tn tntn t tn t tn t

cycloalkadienol) cycloalkadiene)

(5.14)

These aromatic or cycloalkadiene species (CtnH2tn-2t) (Steps Cyclization 1, 2, 2.1, and 3 in

Scheme 5.2) can act as hydrogen donors (denoted as R’H2 in Eqn. 5.2a and Scheme 5.3) for the

intramolecular C=C bond formation (Cycle 2, Scheme 5.2), as discussed in Sections 5.3.2 and

5.3.3. These H donors may also donate their hydrogen atoms to protonated alkenals (CtnH2tn-

2t+2OH+, n=3-5, t=2-4), possibly through the hydride transfer steps described in Section 5.3.2,

163

forming CtnH2tn-2t+2 instead of CtnH2tn-2t (n=3-6, t=2-4). This route, however, is a minor pathway

with the carbon selectivities, defined as CtnH2tn-2t+2 (CtnH2tn-2t+2+ CtnH2tn-2t)−1

, smaller than 15%

and thus will not be addressed here.

Cyclization of long chain (carbon number >5) alkenals such as citral [35-39], citronellal [38, 40],

and aromatic aldehydes (e.g., β-styrylacetaldehyde)[41-44] catalyzed by hydronium ions derived

from HCl and H2SO4 has been studied extensively in the aqueous phase. The reaction is initiated

by protonation of their carbonyl group, followed by an intramolecular electrophilic attack of the

positively charged carbonyl carbon to the C=C double bond, creating an intramolecular C-C

bond and a hydroxyl group and forming cycloalkenol as the product [37, 38]. Sequential

dehydration of the cycloalkenol removes the hydroxyl group and creates the C=C bond, leading

to the formation of a stable aromatic ring. The cycloalkenol can also rearrange via a parallel

reaction that forms cycloketone through keto-enol tautomerization. I hypothesize that the

cyclization of larger alkenal products in alkanal reactions (e.g., 2,4-diethyl-2,4-octadienal from

butanal reactions) occurs via a mechanism similar to those of citral and citronellal [37, 38]. I

probe this reaction using 2,4-heptadienal (C7H10O), because its structure and functional groups

resemble 2,4-diethyl-2,4-octadienal produced from intermolecular C=C bond formation between

butanal and 2-ethyl-2-hexenal (Step 3a, Scheme 5.1).

The reactions of 2,4-heptadienal (C7H10O) on H4SiW12O40 clusters (573 K) led predominantly to

C14H18O (alkenals and isomers), C7H10O (alkyl cycloalkenones), and C7H8 (toluene and a five-

membered ring cycloalkadiene), with carbon selectivities of 63.4 %, 29.3 %, and 4.8 %,

respectively (Fig. S5.6, Appendix). These product distributions suggest that 2,4-heptadienal

(C7H10O) reactions occur via two major pathways: (i) intermolecular C=C bond formation

between two C7H10O reactants, producing larger C14H18O alkenals and (ii) cyclization-

dehydration leading to the formation of cycloalkenones and cyclic hydrocarbon species. Scheme

5.4a shows the proposed mechanism for 2,4-heptadienal cyclization-dehydration, similar to that

for citral and citronellal [37, 38]. The cyclization is initiated by the protonation of 2,4-

heptadienal (Step 1, Scheme 5.4) and the following intramolecular C-C bond coupling (Step 2)

and deprotonation (Step 3) that lead to a five-membered ring (5-MR) alkyl cyclopentadienol. The

cyclopentadienol can either isomerize to form the more stable 5-MR cycloalkenone isomers

(Step 4, carbon selectivity 43 % within all C7 products) or dehydrate to evolve a 5-MR

cycloalkadiene (Step 5, carbon selectivity 3 %). The protonated 2,4-heptadienal could undergo

164

C=C bond shift (Step 1’) followed by cyclization to form a 6-MR cyclohexadienol (Step 2’-3’)

and this route would lead to the formation of six-membered ring (6-MR) cycloalkenones (Step 4’,

carbon selectivity 38 %) and aromatic product, toluene (Step 5’, carbon selectivity 10 %).

Scheme 5.4. (a) Proposed mechanism for acid catalyzed cyclization-dehydration of 2,4-

heptadienal (C7H10O) [the products detected are labeled with carbon selectivities within the C7

product fractions during 2,4-heptadienal reactions on H4SiW12O40 at 573 K, 2,4-heptadienal

165

pressure 0.2 kPa, space velocity= 0.009 +1

Hmol (mol s) , time-on-stream=125 min]; (b)

proposed mechanism for acid catalyzed cyclization-dehydration of 2,4-diethyl-2,4-octadienal

(C12H20O) during butanal reactions [the products detected are labeled with carbon selectivities

within the C12 product fractions during butanal reaction on H4SiW12O40 at 573 K, butanal

pressure 1.1 kPa, space velocity=0.045 +1

Hmol (mol s) , time-on-stream=125 min].

For the secondary cyclization-dehydration of 2,4-diethyl-2,4-octadienal (C12H20O) during

butanal reactions, one could only determine the chemical formulae but not the detailed chemical

structures for most of the products by mass spectrometer. Nevertheless, it is plausible to propose

a cyclization-dehydration mechanism based on the study using 2,4-heptadienal (C7H10O) in

Scheme 5.4a. Scheme 5.4b shows the proposed cyclization-dehydration pathways for 2,4-

diethyl-2,4-octadienal (C12H20O). The cyclization reactions involve carbonyl protonation (Step I),

C=C bond shift (Step I’), intramolecular C-C bond coupling (Step II or II’), and deprotonation

(Step III or III’), evolving 5-MR or 6-MR alkyl cycloalkadienols. These cycloalkadienols can

undergo either isomerization (Step IV or IV’) to form 5-MR or 6-MR cycloalkenone isomers

(C12H20O, carbon selectivity 15.9 % within all C12 products) or dehydration (Step V or V’) to

evolve 5-MR or 6-MR cycloalkadienes (C12H18, carbon selectivity 32.9 % within all C12

products). It was noted that the branched structure of 2,4-diethyl-2,4-octadienal resulted in

diverse cyclization-dehydration products due to the concomitant isomerization reactions. In fact,

there were more than 20 different types of C12H18 products being detected with the highest

selectivities (4.4 %, within all C12 products) towards 1,3,5-triethylbenzene.

These results led to the proposed pathway for alkanal chain growth in Scheme 5.1. Larger

alkenals (CtnH2tn-2t+2O, n=3-6, t=2-3) produced from the sequential, intermolecular C=C bond

formation steps (Steps 2a and 3a) undergo cyclization (Steps 2b and 3b) and evolve

cycloalkenols/cycloalkadienols (CtnH2tn-2t+2O, n=3-6, t=2-3). The cycloalkenols/cycloalkadienols

can either rearrange and desorb as cycloalkanones/cycloalkenones (Steps 2c and 3c) or dehydrate

(Steps 2d and 3d) to evolve unsaturated hydrocarbon species (including cycloalkadienes and

aromatics, CtnH2tn-2t, n=3-6, t=2-3). These pathways led almost exclusively to hydrocarbons with

defined carbon numbers tn, as shown in Figure 5.3, because transalkylation reactions of the

166

aromatic products, which shift the methyl or ethyl groups among the aromatics, remained

kinetically insignificant. As an example, reactions of butanal (C4H8O) on H4SiW12O40 clusters at

573 K produced almost exclusively cycloalkadienes or aromatics with 8, 12, or 16 carbon atoms

(carbon selectivities >99 %, Fig. 5.3b). In contrast, the same reactions on H-form zeolites (H-

ZSM-5[11, 18] and H-Y[18]) produced aromatic products with diverse carbon numbers ranging

from C6 to C16+. The distinct carbon distributions between the H4SiW12O40 clusters and

microporous crystalline materials were caused by effective transalkylation reactions inside the

pores and cages of zeolites, which was not occurred on H4SiW12O40 clusters.

Previous studies have shown that Keggin polyoxometalate clusters (H3PW12O40 and

H4SiW12O40), either dispersed on SiO2[45] or supported on the external surfaces of dealuminated

zeolite Y [46], could catalyze the transalkylation of alkyl benzenes (trimethylbenzene and

toluene), leading to C6 to C10+ aromatics at moderate temperatures (573-723 K). Reactions of

tetralin (C10H12) in the absence of alkanals could provide the reactivities of aromatic

transalkylation. In the absence of an alkanal, tetralin reactions on H4SiW12O40 formed not only

C10 naphthalene as the dehydrogenation product, but also C6-C13 aromatics via transalkylation

reactions. The carbon selectivities towards the transalkylation products, (1−C10/13

6C

m

mm

),

were 7.2 % and the transalkylation rate was 0.40 +1

tetralin Hmmol (mol s) at 573 K [space

velocity=5.6 +1

tetralin Hmmol (mol s) ]. These selectivity values were similar to tetralin reactions

on H-Y zeolite (Si/Al=15), where the transalkylation selectivity ( 1−C10/13

6C

m

mm

) was 20 %

and transalkylation rate was 0.11 +1

tetralin Hmmol (mol s) at 573 K [space velocity=0.57

+1

tetralin Hmmol (mol s) ]. These selectivities contradicted those on H4SiW12O40 clusters after

butanal catalysis at 573 K for >155 min, which showed no transalkylation reactivity of tetralin at

573 K. These results, together with the lack of transalkylation of aromatic products (C8, C12, and

C16) detected during butanal reactions at 573 K (Fig. 5.3b), indicate that butanal adsorption

inhibits the transalkylation activity on H4SiW12O40 clusters. The lack of detectable

transalkylation reactivities during alkanal reactions on H4SiW12O40 is likely caused by the loss of

strong H+ sites resulted from their binding to heavier products (e.g., cokes), as confirmed from

temperature programmed desorption of pyridine carried out on H4SiW12O40 clusters after steady-

state reactions (See Fig. S5.5, Appendix). Similar conclusions have also been shown previously

167

for transalkylation of alkyl aromatics on a series of HNa-Y and H-USY zeolites, which indicated

that only strong Brønsted acid sites with NH3 desorption temperature above 623 K were active in

transalkylation reactions [47].

5.3.5. Effects of alkanal chain length on its deoxygenation rates and selectivities on H4SiW12O40 clusters

Figure 5.6a shows the rate constants ( Inter,effk , Intra,effk , Dehy,bi,effk , and Dehy,mono,effk in Equations

5.5, 5.8, 5.11, and 5.12, respectively) for the primary pathways (Cycles 1, 2, 3, and 3.1, in

Scheme 5.2) of C3-C6 n-alkanal reactions on H4SiW12O40 at 573 K and the selectivities for

secondary cyclization-dehydration reactions (Cyclization 2, Scheme 5.2), 3Cycli-dehy,C n

, defined as

the site-time-yield of C3n aromatics (3C aromn

r ) divided by that of all C3n products, including

oxygenates and aromatics (3C overalln

r ):

3

3

3

C arom

Cycli-dehy,C

C overall

n

n

n

r

r (5.15)

Figure 5.6. (a) Rate constants for intermolecular C=C bond formation ( Inter,effk , , Cycle 1 in

Scheme 5.2), intramolecular C=C bond formation ( Intra,effk , ●, Cycle 2), and alkanal

isomerization-dehydration via bimolecular pathway ( Dehy,bi,effk , , Cycle 3) and monomolecular

168

pathway ( Dehy,mono,effk , ▲, Cycle 3.1) and cyclization-dehydration selectivity of C3n alkenal

(3Cycli-dehy,C n

, ♦, Cyclization 2 and 2.1 in Scheme 5.2, Eqn. 5.15) during alkanal (CnH2nO, n=3-6)

deoxygenation on H4SiW12O40 clusters at 573 K as a function of reactant carbon number

[CnH2nO=1.1-10 kPa, 0.045-0.44 +1

alkanal Hmol (mol s) , alkanal conversion=14-17 %, 26-31 %,

45-47%, and 68-72 % for propanal, butanal, pentanal and hexanal, respectively]; (b) first-order

rate constants for intramolecular C=C formation ( Intra,effk ) for C3-C6 alkanals (CnH2nO, n=3-6) on

H4SiW12O40 clusters at 573 K as a function of the hydride ion affinity difference ( HIA )

between the carbenium ions of H-donor (R’H+) and the protonated alkanal (CnH2nOH

+) [ HIA =

+R'HHIA − +

2C H OHn nHIA , where R’H

+=C10H11

+, as tetralin (C10H12) is used as the representative H-

donor to estimate HIA values].

Next, I decompose the effective rate constants ( Inter,effk , Intra,effk , Dehy,bi,effk , and Dehy,mono,effk ) to

elementary rate and equilibrium constants and connect these reactivity trends to the

thermodynamic properties of Brønsted site and reactants. According to Equations 5.5 and 5.11,

the effective rate constant ratio for intermolecular C=C bond formation ( Inter,effk ) to

isomerization-dehydration ( Dehy,bi,effk ) equals the rate constant ratio for aldol condensation ( C-Ck ,

Step R1.2) to alkanal isomerization ( iso,bik , Step R3.1a):

Inter,eff C-C

Dehy,bi,eff iso,bi

k k

k k (5.16)

As shown in Figure 5.6a, during steady-state reaction, Inter,effk is much higher than Dehy,bi,effk

with Inter,eff Dehy,bi,eff

1k k

ratios of 114, 173, and 217 for butanal, pentanal, and hexanal,

respectively, thus C-Ck is much larger than iso,bik . Therefore, the effective rate constants for the

intermolecular C=C formation (Eqn. 5.5) and the alkanal dehydration via bimolecular pathway

(Eqn. 5.11) can be simplified further to:

169

tautC-C AAPInter,eff

-AAP C-C

k k Kk

k k

(5.17)

tautAAPiso,bi

Dehy,bi,eff

-AAP C-C

k k Kk

k k

(5.18)

Equation 5.17 indicates that Inter,effk depends on the rate constants for alkanal-alkenol pair

(AAP*) formation ( APPk , Step R1.1), its reverse reaction ( -APPk , Step R1.1’), and aldol

condensation ( C-Ck , Step R1.2), as well as by the equilibrium constant for keto-enol

tautomerization ( tautK , Step G1).

As the alkanal reactant size increases from C3 to C6, the effective rate constant Inter,effk increases

(as shown in Fig. 5.6a), reflecting the increase in the values of either APPk , C-Ck , and/or tautK .

Little information on tautK values is available for the keto-enol tautomerization of gaseous C3-C6

n-alkanals. However, it is known that the equilibrium constant for keto-enol tautomerization

( tautK ) of isobutyraldehyde in the aqueous phase is much higher than that of acetaldehyde

(1.28×10−4

vs. 5.89 ×10−7

, at 298 K) [48-50], which suggests that larger alkyl substituent favors

the enol formation and thus exhibits higher tautK values. It is because the σ-π hyperconjugation

between the alkyl substituent and the enol C=C bond delocalizes electrons of the alkyl group

onto the C=C bond and stabilizes the enol;[48, 51, 52] larger alkyl substituent promote the enol

stabilization to a greater extent.

I expect that the larger alkanals favor the formation of the adsorbed alkanal-alkenol pair (AAP*,

Step R1.1) and the bimolecular transition state for aldol condensation [TS(C-C)*, Step R1.2] on

H4SiW12O40 clusters, because of their stronger van der Waals interactions with the catalyst

surfaces [53]. As an example, van der Waals interactions increase with the carbon number in n-

alkanes, causing the heats of n-alkane adsorption to increase by 1.5-2 kJ mol−1

for each

additional C atom, when adsorbing them on mesoporous silica structures [54]. Therefore, I

expect the larger alkanal to exhibit a higher rate constant for AAP* formation ( APPk , Step R1.1).

In the aldol condensation step (Step R1.2), both the reactant state (AAP*) and transition state

[TS(C-C)*] contain the same carbon number and therefore the extent of stabilization remains the

170

same. For this reason, I expect that the activation barrier for TS(C-C)* formation and the related

C-Ck remain insensitive to the alkanal reactant size. Thus, the higher effective rate constant

Inter,effk for larger alkanals must reflect their higher tautK and APPk values, which correspond to

the higher stability of enol tautomer and more abundant AAP* intermediates, respectively, than

the smaller alkanals.

The effective rate constants for alkanal dehydration via the monomolecular pathway (Steps

R3.1b-R3.2b), Dehy,mono,effk , remain relatively stable for C4-C6 alkanals, as shown in Figure 5.6a,

thus the alkanal isomerization step ( iso,monok , Step R3.1b) is not sensitive to the alkanal reactant

size. It is plausible that the rate constant for alkanal isomerization via the bimolecular pathway

( iso,bik , Step R3.1a) is also not sensitive to the reactant size. As a result, the larger alkanals with

higher tautK and APPk values exhibit a higher effective rate constant Dehy,bi,effk according to

Equation 5.18, as shown in Figure 5.6a.

In contrast to the rate constant trends for the intermolecular C=C bond formation and

isomerization-dehydration reactions ( Inter,effk and Dehy,bi,effk ), the effective rate constant for

intramolecular C=C bond formation ( Intra,effk ) decreases with increasing alkanal’s carbon

number. As shown in Equations 5.7 and 5.8, the effective rate constant Intra,effk reflects the

elementary rate constant for the hydride transfer step ( HTk , Step R2.1, Scheme 5.3). During the

hydride transfer, H-donors (R’H2) donate their hydride ions to protonated alkanals (CnH2nOH+)

and convert into carbenium ions (R’H+) (Step R2.2), because CnH2nOH

+ has a higher hydride ion

affinity (HIA) than R’H+. The hydride ion affinity difference (∆HIA) between the carbenium ion

of the H-donor (R’H+, +R'H

HIA ) and the protonated alkanal as the H-acceptor (CnH2nOH+,

+2C H OHn n

HIA ) dictates the hydride transfer reactivity:

HIA = +R'HHIA − +

2C H OHn nHIA (5.19)

The H-donor-acceptor pairs with more negative ∆𝐻𝐼𝐴 values exhibited higher hydride transfer

rates. Because the alkyl tetralins in the aromatic product fractions are the major H-donors for

171

alkanal transfer hydrogenation [18], I use tetralin (C10H12) as the representative H-donor

(R’H+=C10H11

+, +R'H

HIA =934.1 kJ mol−1

[18]) to estimate the ∆HIA values for different alkanals.

As the carbon number (n) of alkanal increased from three to six, its +2C H OHn n

HIA value decreased

from 956.6 kJ mol−1

to 941.1 kJ mol−1

and the ∆HIA increased from −22.5 kJ mol−1

to −7.0 kJ

mol−1

. The less negative ∆HIA values led the Intra,effk to concomitantly decrease from 11.5 to

0.48 +1

Hmmol (mol s kPa) (573 K, Fig. 5.6b). This direct correlation between the reactivities

and the ∆HIA has also been demonstrated previously on H-FAU zeolites [18].

The reactivity of the secondary cyclization-dehydration reactions also increased with the alkanal

size, as indicated by the higher molar percentages of C2n cycloalkadienes and C3n aromatics in

the C2n and C3n product fractions (Fig. 5.3) and higher cyclization-dehydration selectivities

towards C3n alkenal (3Cycli-dehy,C n

, Eqn. 5.15, Fig. 5.6a) for the larger alkanals. The cyclization-

dehydration pathway requires the electrophilic attack of the carbonyl carbon onto the C=C

double bond, a step promoted by an alkyl substitution at the C=C position, because the

substitution leads to higher electron densities at the C=C bond [42]. As the chain length of the

alkanal increases, the larger alkyl group at the C=C position (–R group, as shown in Scheme 5.1)

affords more effective electron donation and thus results in larger cyclization-dehydration rates.

5.4. Conclusion

Kinetic measurements and acid site titration lead to a proposed reaction network with parallel

and sequential catalytic cycles for the deoxygenation of light alkanals (CnH2nO, n=3-6) catalyzed

by the Brønsted acid sites (H+) on tungstosilicic acid clusters (H4SiW12O40). Alkanal

deoxygenation proceeds via three primary pathways of: (1) intermolecular C=C bond formation,

which couples two alkanal molecules via kinetically-relevant aldol-condensation step followed

by a rapid dehydration step, that evolves a larger alkenal (C2nH4n-2O), (2) intramolecular C=C

bond formation, which converts alkanals directly to n-alkenes (CnH2n), via kinetically-relevant

hydride ion transfer from H-donating agents to protonated alkanals, followed by dehydration,

and (3) isomerization-dehydration, during which the alkanals first isomerize to form allylic

alcohols then rapidly dehydrate to produce n-dienes (CnH2n-2). In the catalytic cycle of Pathway

172

(1), a series of sequential intermolecular C=C bond formation events adds additional alkanal

units onto the carbon chain, thus producing larger alkenals (C3nH6n-4O and C4nH8n-6O). These

larger alkenal species can undergo cyclization-dehydration reactions, leading to cyclic

hydrocarbons including cycloalkadiene (C2nH4n-4) and aromatic species (C3nH6n-6 or C4nH8n-8).

The catalytic pathways are kinetically coupled together, because cyclic hydrocarbons produced

from the sequential reactions of Pathway (1) act as the hydrogen donors for Pathway (2) and

Pathways (2) and (3) share the co-adsorbed alkanal-alkenol pairs as the common reaction

intermediates.

The molecular size of alkanals affects their thermochemical properties and in turn influences the

stabilities of the transition states and reaction intermediates in the kinetically-relevant steps of

the different pathways. These effects lead to contrasting reactivity trends for the various reaction

pathways, thus resulting in different selectivities across the alkanal family. The rate constants for

Pathways (1) and for the bimolecular route of Pathway (3) both increase with alkanal size,

apparently because both reactions require enol tautomers, which are more stable for larger

alkanals, for the formation of bimolecular alkanal-alkenol pairs as the reaction intermediates.

Alkanal size does not affect the rate constants for the monomolecular route of Pathway (3),

because protonated alkanal monomers remain as the most abundant surface intermediates. In

contrast, the rate constants for Pathway (2) decrease with increasing alkanal size, because larger

alkanals exhibit lower hydride ion affinities and thus are less effective towards hydride ion

abstraction. The reactivities of the secondary alkenal cyclization-dehydration reactions increase

with molecular size, because larger alkyl substitution at the C=C position of the alkenals

increases the electron density of the C=C bond and thus promotes the intramolecular

electrophilic attack of the carbonyl group onto the C=C bond to initiate the cyclization-

dehydration reactions. This mechanistic knowledge on the tandem catalytic cycles and their

kinetic and thermodynamic requirements provide the framework for rationalizing and then

predicting the site-time-yields for larger oxygenate and hydrocarbon during alkanal

deoxygenation turnovers.

173

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5.6. Appendix

5.6.1. Characterizations of Brønsted and Lewis acid sites on H4SiW12O40

The Brønsted and Lewis acid sites on H4SiW12O40/SiO2 catalysts were characterized by pyridine

titration and by the infrared spectroscopic study of pyridine adsorption. Table S5.1 summarizes

the amounts of Brønsted and Lewis acid sites on these catalysts. The total amount of acid sites

was determined based on the pyridine uptake during pyridine titration at 473 K. The percentage

of Brønsted and Lewis acid sites were determined based on the infrared absorption bands at 1545

and 1455 cm−1

, respectively, measured during pyridine adsorption on the catalysts at 473 K [55].

Table S5.1. The amounts of Brønsted and Lewis acid sites on H4SiW12O40 catalysts

H4SiW12O40

Total acid sites (μmol gcat.−1

)a 169±6

Brønsted acid percentageb 93.7 %

Lewis acid percentageb 6.3 %

Brønsted acid sites (μmol gcat.−1

) 159±6

Lewis acid sites (μmol gcat.−1

) 11±1

aThe amounts of total acid sites were determined by pyridine titration at 473 K; bThe percentage of Brønsted and Lewis acid sites were determined from the infrared absorption spectra taken during pyridine

adsorption at 473 K [55].

178

5.6.2. Stability of polyoxometalate at high temperature

The thermal stability of the polyoxometalate clusters was examined using H3PW12O40. The

H3PW12O40/SiO2 catalysts (loading amount 0.1 3 12 40 2

1H PW O SiO

mmol g ) were prepared with the

same method stated in this manuscript. The H3PW12O40/SiO2 samples were pretreated at 473,

573, and 673 K (in 50 cm3 min

−1 He for 3 h), respectively, before the pyridine titration at 473 K.

Figure S5.1 shows the amount of H+ sites on these pretreated H3PW12O40 clusters determined by

pyridine titration. The thermal treatment at 473 K does not affect the H+ site density on

H3PW12O40 (3.06 molH+ molcluster−1

, identical to the theoretical value). The treatments at 573 K

and 673 K decrease H+ site density by only 6 % and 17 %, respectively (to 2.89 and 2.53

molH+ molcluster−1

, respectively), probably due to H3PW12O40 dehydration, indicating that the

polyoxometalate cluster is stable at 573 K. In contrast, butanal reactions at 573 K lead the

accessible H+ sites on H4SiW12O40 to decrease by as much as ~40 % (from 2.56 to 1.54

molH+ molcluster−1

, Fig. 5.1b), which must be caused by coke deposition instead of the

dehydration of polyoxometalate.

Figure S5.1. H+ site densities on H3PW12O40 clusters after thermal treatment under flowing He at

different temperatures (473-677 K) for 180 min (loading amount 0.13 12 40 2

1H PW O SiO

mmol g ).

400 500 600 7000

1

2

3

4

Acid

site

de

nsity (

mo

l H+ m

ol c

luste

r-1)

Pretreatment temperature (K)

179

5.6.3. Coke formation during butanal reactions on H4SiW12O40

On H4SiW12O40, butanals undergo intermolecular C=C bond formation and sequential

cyclization-dehydration, forming aromatic products (e.g. triethylbenzene and alkyl tetralins) at

573 K. On fresh H4SiW12O40, these aromatic species further undergo dehydrogenation, leading to

the formation of cokes. Figure S5.2 shows the differential spectra of H4SiW12O40 before and after

butanal reactions at 573 K for 5 min. The band at 1580 cm−1

is ascribed to the stretching

vibration of the C=C bond in aromatic rings. The appearance of this C=C vibrational band

without the concomitant appearance of the C-H stretching band indicates the formation of cokes.

These features are similar with those of pure graphene sheet [56].

Figure S5.2. The differential spectra of H4SiW12O40 before and after butanal reactions at 573 K

for 5 min.

The coke formation during butanal reaction was further confirmed by temperature program

oxidation (TPO) of the spent H4SiW12O40 catalysts (after butanal reactions in 0.5 kPa butanal, at

573 K, for 8 h) using TG-DSC (thermogravimetry-differential scanning calorimetry methods).

Figure S5.3 shows the weight, differential weight loss, and heat flow profiles for the TPO of the

spent catalyst samples (initial weight m0=15.3 mg, heating rate=1 K min−1

, in 5 % O2/He, flow

rate=10 cm3 min

−1). There are three ranges of weight loss during the TPO:

4000 3500 3000 2500 2000 1500 1000

0.00

0.02

0.04

0.06

0.08

0.10

0.12

3740 cm-1A

bso

rba

nce

(a

.u.)

Wavenumber (cm-1)

1580 cm-1

180

(i) The weight loss below 600 K (∆m<600K=1.72 mg) is likely ascribed to the desorption of water

and some organic species (e.g., larger alkenals and heavy aromatic products formed during

butanal reactions);

(ii) The weight loss in the temperature range of 600-800 K (∆m600-800K=0.68 mg) is ascribed to

the oxidation of coke species [57], as indicated by the exothermicity of the oxidation reaction;

(iii) The slight weight loss above 800 K (∆m>800K=0.15 mg) is probably caused by the

dehydration of polyoxometalate clusters [58].

I used the weight loss data between 600-800 K, together with Equation (S5.1), to obtain the

amount of coke deposited (wcoke) on the catalysts.

𝑤coke =∆𝑚600−800K

𝑚0−∆𝑚<600K−∆𝑚600−800K (S5.1)

In Equation S5.1, 𝑚0 is the initial weight of spent catalyst sample; ∆𝑚<600K and ∆𝑚600−800K are

the weight losses in the temperature ranges of <600 K and 600-800 K, respectively. The amount

of coke deposited is determined to be 5.2 wt.%.

Figure S5.3. Weight, differential weight loss, and heat flow profiles during the temperature

programmed oxidation (TPO) of spent H4SiW12O40/SiO2 catalysts (0.075 4 12 40 2

1H SiW O SiO

mmol g )

400 600 800 1000-5

0

5

10

15

Diffe

ren

tia

l w

eig

ht

loss (g

s-1

)

We

igh

t (m

g)

Sample Temperature (K)

Weight

Differential weight loss

-20

-10

0

10

20

He

at

flo

w (

mW

)

Heat flow

181

after butanal reactions (0.5 kPa butanal) at 573 K for 8 h (initial sample weight=15.3 mg, in 5%

O2/He, 10 cm3 min

−1, heating rate =5 K min

−1).

5.6.4. Effects of extra-framework alumina on the butanal reaction on H-MFI zeolite

The H-MFI zeolite (denoted as H-MFI, in Table S5.2 and Fig. S5.4), was treated by ammonium

hexafluorosilicate (NH4)2SiF6 to remove the extra-framework alumina [denoted as H-

MFI(AHFS), in Table S5.2 and Fig. S5.4]. The FTIR spectroscopic study of pyridine adsorption

was performed to quantify the amount of Brønsted and Lewis sites on these two zeolite samples

as listed in Table S5.2. The lower Lewis acid site density on H-MFI(AHFS) indicates less extra-

framework alumina on this sample. Other than the Lewis acidity, the extra-framework alumina

also contains basic sites according to the FTIR study of the CO2 and boric acid trimethyl ester

adsorption on γ-alumina [59]. The quantum chemical study of alumina indicates that O atom

bridging to two Al atoms exhibits basicity [34]. Although I do not directly measure the amount

of basic sites on the zeolite samples, it is plausible that H-MFI(AHFS), which has less extra-

framework alumina, has less basic sites than H-MFI.

Figure S5.4 shows the rates for intermolecular C=C bond formation ( Interr ), intramolecular C=C

bond formation ( Intrar ), and isomerization-dehydration ( Dehyr ) during butanal reactions at 573 K

on H-MFI and H-MFI(AHFS) as a function of time-on-stream. Despite deactivation due to heavy

product deposition, H-MFI and H-MFI(AHFS) both show similar Interr and Intrar (Fig. S5.4a and

S5.4b), indicating that the extra-framework alumina has little effects on the pathways of inter-

and intramolecular C=C bond formation. In contrast, 𝑟Dehy on H-MFI(AHFS) is about half of

that on H-MFI (Fig. S5.4c), an indication that the extra-framework alumina promotes the

isomerization-dehydration reaction. Because the isomerization of alkanal to allylic alcohol

requires a basic site as proton abstractor to complete the shift of C=O bond to C=C bond, as

shown by Steps R3.1b-R3.2b in Scheme 5.3, and also proposed in previous study on 2-

methylbutanal dehydration [17], it is likely that the extra-framework alumina on the zeolite

provides the basic sites rather than Lewis acid sites to catalyze the butanal isomerization.

182

Table S5.2. Amount of Brønsted and Lewis acid sites on H-MFI zeolites (Si/Al=40) and the

ammonium hexafluorosilicate treated H-MFI zeolites [H-MFI(AHFS)] (measured by FTIR study

of pyridine adsorption at 448 K)

H-MFI H-MFI(AHFS)

Brønsted acid site density (μmol gcat.−1

) 380±20 340±20

Lewis acid site density (μmol gcat.−1

) 62±10 22±3

Figure S5.4. Rates for (a) intermolecular C=C bond formation ( Interr ), (b) intramolecular C=C

bond formation ( Intrar ), and (c) isomerization-dehydration ( Dehyr ) during butanal reaction on H-

MFI zeolite (Si/Al=40) and the ammonium hexafluorosilicate treated H-MFI zeolite [H-

MFI(AHFS)] at 573 K as a function of time-on-stream [butanal=4 kPa, space velocity=1.5×10−5

mol (gcat. s)−1

].

0 50 100 1500

5

10

15

r Inte

r (1

0-7

mo

l (g

cat. s

)-1)


H-MFI

H-MFI(AHFS)

0 50 100 1500

5

10

15

r Intr

a (

10

-7 m

ol (

gca

t. s

)-1)


H-MFI

H-MFI(AHFS)

0 50 100 1500.0

0.5

1.0

1.5

r De

hy (

10

-7 m

ol (

gca

t. s

)-1)

Time-on-strem (min)

H-MFI

H-MFI(AHFS)

rInter rIntra

rDehy

(a) (b) (c)

183

5.6.5. Temperature programmed desorption (TPD) of pyridine on H4SiW12O40

Temperature programed desorption (TPD) of pyridine was performed to measure the strength

and the amount of H+ sites on the fresh and spent H4SiW12O40 catalysts. Three desorption peaks

centered at 560 K, 660 K, and 720 K, respectively, were observed during pyridine-TPD on the

fresh H4SiW12O40 catalyst as shown in Figure S5.5a. These desorption peaks correspond to three

types of H+ sites with different acid strengths. After 12 h of butanal reactions at 573 K, only two

desorption peaks at 540 K and 640 K were observed (Fig. S5.5b) from the spent H4SiW12O40

catalyst, indicating the loss of strong H+ sites (with pyridine desorption peak of 720 K) due to the

heavy product deposition.

Figure S5.5. Profiles of pyridine-TPD for (a) fresh H4SiW12O40 catalysts and (b) spent

H4SiW12O40 catalysts after 12 h of butanal reactions at 573 K (pyridine adsorption temperature

473 K, heating rate 1 K min−1

).

5.6.6. Reactions of 2,4-heptadienal on H4SiW12O40

The pathways for secondary cyclization and dehydration were probed with 2,4-heptadienal

(C7H10O), because its structures and functional groups resemble the 2,4-diethyl-2,4-octadienal

500 600 700 8000.000

0.005

0.010

0.015

Pyri

din

e d

eso

rptio

n r

ate

(mo

l (m

ol H

4S

iW1

2O

40 m

in)-1

)

Temperature (K)

500 600 700 8000.000

0.005

0.010

0.015P

yri

din

e d

eso

rptio

n r

ate

(mo

l (m

ol H

4S

iW1

2O

40 m

in)-1

)

Temperature (K)

560 K

660 K 720 K(a) (b)

540 K

640 K

184

produced from intermolecular C=C bond formation between the butanal and 2-ethyl-2-hexenal

(Step 3a, Scheme 5.1 of Sec. 5.3.1). Figure S5.6 shows the carbon product distributions of 2,4-

heptadienal reactions on H4SiW12O40 at 573 K. The reactions form predominantly condensation

products (C14H18O, 64 % carbon selectivity) and cyclization products (C7H10O, e.g. alkyl

cyclohexenones and alkyl cyclopentenones, total carbon selectivity of 29 %). The

cycloalkenones are the products of the cyclization of heptadienals that can further undergo

dehydration and rearrangement to form C7H8 hydrocarbons (e.g. toluene and the five-membered

ring isomer, 4.8 % total carbon selectivity).

Figure S5.6. Carbon selectivities of the products during 2,4-heptadienal reactions on

H4SiW12O40 clusters at 573 K [2,4-heptadienal pressure=0.2 kPa, space velocity=0.009

+1

Hmol (mol s) , time-on-stream=125 min]

0

20

40

60

80

C7H

6O

Be

nza

lde

hyd

e

C14H

18O

C7H

10O

iso

me

rs

e.g

. cyclo

pe

nte

no

ne

s

C7H

10

C7H

8 iso

me

r

Ca

rbo

n s

ele

ctivitie

s (

%)

C7H

8

To

lue

ne

185

Chapter 6 Summary of Alkanal Deoxygenation on Solid Brønsted Acid

Sites and Perspective for Bio-oil Upgrading

6.1. Summary of catalytic pathways for alkanal deoxygenation

Alkanals are deoxygenated on solid Brønsted acid catalysts (e.g. acidic zeolites and

polyoxometalate clusters) at moderate temperatures (473-673 K), forming multiple products

including heavy alkenals, aromatics, light alkenes, and alkadienes. The work of this thesis

established the three primary deoxygenation routes: (i) intermolecular C=C bond formation,

which occurs between two alkanal molecules (CnH2nO) via the aldol condensation-dehydration

reaction, producing larger alkenals (C2nH4n-2O), (ii) intramolecular C=C bond formation, which

proceeds via transfer hydrogenation of one alkanal molecule (CnH2nO), followed by dehydration

evolving light alkene (CnH2n) while preserving the carbon backbone, (iii) direct-dehydration

pathway, which converts alkanal (CnH2nO) into light alkadiene (CnH2n-2). This work gains the

mechanistic insight into these three alkanal deoxygenation pathways and probes their kinetic

requirements and kinetic couplings.

6.1.1 Intermolecular C=C bond formation

The pathway of intermolecular C=C bond formation is initiated by the adsorption and

protonation of alkanal molecules (CnH2nO) on the Brønsted acid sites (H+). In the gas phase,

there is a keto-enol tautomerization equilibrium that forms the alkenol tautomer (CnH2n-1OH) of

the alkanal. The nucleophilic attack of the alkenol (CnH2n-1OH) onto the protonated alkanal

(CnH2nOH+) leads to the aldol condensation, creating an intermolecular C-C bond and forming an

aldol (C2nH4nO2, Eqn. 6.1a). The rapid dehydration of the aldol produces an alkenal (C2nH4n-2O,

Eqn. 6.1b) and completes the catalytic cycle [1]. This intermolecular C=C bond formation route

can add more alkanal units and augment the chain length of the alkenal products (e.g. C3nH6n-4O

and C4nH8n-6O, Eqns. 6.1c-6.1d). The sequential cyclization-dehydration of the alkenal species

strips their last oxygen atoms and converts them into unsaturated cyclic hydrocarbon products,

including cycloalkadienes (e.g., C2nH4n‒4, Eqn. 6.1e) and aromatics (e.g., C3nH6n-6 and C4nH8n-8,

Eqns. 6.1f-6.1g). This chain growth route is an effective route to produce alkenal species that are

important intermediates for the production of fine chemicals used for perfumes [2], flavors [3, 4],

186

and pharmaceutical products [5]; it can also lead to heavy aromatic species used for blending

with hydrocarbon fuels (e.g., diesel and jet fuel).

2 CnH2nO→C2nH4nO2 + H2O (6.1 a)

C2nH4nO2→C2nH4n‒2OH + H2O (6.1 b)

C2nH4n‒2O + CnH2nO→C3nH6n‒4O + H2O (6.1 c)

C3nH6n‒4O + CnH2nO→C4nH8n‒6O + H2O (6.1 d)

C2nH4n‒2O→C2nH4n‒4 + H2O (6.1 e)

C3nH6n‒4O→C3nH6n‒6O + H2O (6.1 f)

C4nH8n‒6O→C4nH8n‒8O + H2O (6.1 g)

The reactivity of the primary intermolecular C=C bond formation is sensitive to the molecular

sizes of the alkanal reactants. Typically, the larger alkanals form more alkenols via keto-enol

tautomerization equilibrium and, as a result, have higher frequencies of collision between

alkenols and protonated alkanals, leading to higher turnover rates of intermolecular C=C bond

formation.

The rates and selectivities of the intermolecular C=C bond formation and the sequential

cyclization-dehydration reactions can be controlled by varying the local confinement around the

H+ sites. The turnover rates for aldol condensation are much higher (one order of magnitude

higher) on the H+ sites dispersed on the polyoxometalate clusters (e.g. H4SiW12O40) than those

confined within microporous zeolites (e.g. H-MFI and H-FAU) [6]. However, the secondary

cyclization-dehydration of the alkenals, which forms aromatic products, preferentially occurs

within confined catalytic environment of zeolites. These kinetic consequences of the local H+ site

structure allow the reactivity and selectivity tuning.

6.1.2. Intramolecular C=C bond formation

The pathway of intramolecular C=C bond formation also takes place on the protonated alkanal

(CnH2nOH+) via a kinetically-relevant transfer hydrogenation step that converts alkanal into an

alcohol (CnH2n+1OH, Eqn. 6.2a), which then rapidly dehydrates to form an alkene (CnH2n, Eqn.

187

6.2b). Without adding any external hydrogen source, the cyclic hydrocarbon products of the

aldol condensation-cyclization routes act as the hydrogen donors (H-donor, denoted as RDH2) [1,

6]. The kinetically-relevant transfer hydrogenation step proceeds via a hydride ion transfer

mechanism, during which the H-donor (RDH2) donates a hydride ion, which attaches onto the

carbonyl carbon of the protonated alkanal (CnH2nOH+) to form a concerted, bimolecular

transition state,[CnH2nOH+∙∙∙H‒∙∙∙RDH

+]

‡. This transition state decomposes to release an alcohol

(CnH2n+1OH) and the carbenium ion of the H-donor (RDH+). The alcohol rapidly dehydrates to

form an alkene (CnH2n), while the carbenium ion (RDH+) donates a proton back onto the catalyst

surface to regenerate the H+ site. This transfer hydrogenation route converts alkanals into their

corresponding alkenes, which can be used as fuels and important raw materials for the

production of polymers and other valuable chemicals (e.g., alcohols); this reaction occurs under

mild conditions (473-673 K and atmospheric pressure), in comparison to alkanal

hydrodeoxygenation with high-pressure H2 (20-100 bar).

CnH2nO + RDH2→CnH2n+1OH + RD (RDH2 represents the H-donor) (6.2a)

CnH2n+1OH→CnH2n + H2O (6.2b)

The rate for the transfer hydrogenation is dictated by the hydride ion affinity difference

(∆𝐻𝐼𝐴RDH+−C𝑛H2𝑛OH+) between the protonated alkanal (CnH2nOH+) and the carbenium ion of the

H-donor (RDH+) [6, 7], defined in Equation 6.3:

∆𝐻𝐼𝐴RDH+−C𝑛H2𝑛OH+= 𝐻𝐼𝐴RDH+ − 𝐻𝐼𝐴C𝑛H2𝑛OH+ (6.3)

where 𝐻𝐼𝐴RDH+ and 𝐻𝐼𝐴C𝑛H2𝑛OH+ are the hydride ion affinities for the carbenium ion of the H-

donor (RDH+) and the protonated alkanal (CnH2nOH

+), respectively. Only the H-donor-acceptor

pairs with negative hydride ion affinity differences (∆𝐻𝐼𝐴RDH+−C𝑛H2𝑛OH+< 0 kJ mol‒1

) can carry

out the transfer hydrogenation reaction; more negative ∆𝐻𝐼𝐴RDH+−C𝑛H2𝑛OH+ would lead to a

lower activation barrier and higher rate for transfer hydrogenation. Within the alkanal family, a

smaller alkanal molecule that has a higher hydride ion affinity (𝐻𝐼𝐴C𝑛H2𝑛OH+) exhibits a higher

tendency to undergo transfer hydrogenation.

The H+ site local environment also influence the reactivity of alkanal transfer hydrogenation [6].

The partially confined structure around H+ sites can solvate and stabilize the bimolecular

188

transition state ([CnH2nOH+∙∙∙H‒∙∙∙RDH

+]

‡) and promote the transfer hydrogenation. As a result,

the hydride transfer between the C3-C6 n-alkanals (kinetic diameter 4.7-6.2 Å [8]) and

hydrocarbon species (e.g., cyclohexadiene and tetralin, kinetic diameter 5.9-6.2 Å [9, 10]) is

more effective when occurs in the supercages of H-FAU zeolite (diameter 11.8 Å) than in the

channels of H-MFI zeolite (diameter 5.6 Å) and in silica mesoporous structure of

H4SiW12O40/SiO2 catalyst (pore size > 150 Å). This kinetic requirement for hydride transfer

allows us to control the reactivity of alkanal transfer hydrogenation by tuning the catalyst

microstructure .

6.1.3. Direct alkanal dehydration

Beside the aldol condensation and alkanal transfer hydrogenation, H+ sites can also catalyze the

isomerization of alkanals (CnH2nO) to form allylic alcohols (CnH2n-1OH, Eqn. 6.4a) followed by

rapid dehydration (Eqn. 6.4b) to evolve alkadienes (CnH2n-2) [7]. This direct alkanal dehydration

route is potentially important for the production of alkadienes as raw materials for polymer

synthesis [11-13]. The kinetically-relevant alkanal isomerization step can take place via a

monomolecular pathway on a protonated alkanal or via a bimolecular pathway on a co-adsorbed

alkanal-alkenol pair; both pathways require an acid-base site pair.

CnH2nO→CnH2n‒1OH (6.4a)

CnH2n‒1OH→CnH2n‒2 + H2O (6.4b)

The route of direct alkanal dehydration is kinetically coupled with the route of alkanal

intermolecular C=C bond formation, via a shared alkanal-alkenol intermediate. The rates of both

routes increase with the alkanal molecular size, because larger alkanals form more stable alkenol

tautomers and thus favor the formation of alkanal-alkenol intermediate.

The kinetically-relevant alkanal isomerization step requires both the H+ site and the conjugated

basic site (the lattice oxygen) on the catalyst surface. Increasing the amount of basic sites (e.g.

extra-framework alumina on zeolites) can promote the rate of alkadiene formation via direct

alkanal dehydration.

189

6.2. Perspective on the study of catalytic deoxygenation of bio-oil on solid Brønsted acid catalysts

Alkanal is only one of the major components in bio-oil. Current separation technologies are

unable to completely extract individual components from the complex composition of bio-oil

[14]. Therefore, the catalytic deoxygenation of the components other than alkanals is also of

great interest for the studies on the catalytic upgrading of bio-oil.

Ketones and phenolic species are two other major components in bio-oil. The deoxygenation

route of transfer hydrogenation (intramolecular C=C bond formation) for alkanals is also

applicable for the ketones and phenolic species, because the ketones also contain a carbonyl

group (C=O) and phenolic species also have equilibrated keto tautomers (with C=O). The

Brønsted acid sites can catalyze the transfer hydrogenation of the C=O bonds of ketones and

phenolic species in the presence of hydride donors, followed by dehydration, to eject the oxygen

heteroatoms. For example, the larger pore H-FAU zeolites are active in catalyzing the transfer

hydrogenation of 1-tetralone to naphthalene [15] and lignin-derived phenols to alkyl benzenes

[16], with tetralin or substituted tetralins as the H-donors. Therefore, mechanistic and kinetic

investigations into the transfer hydrogenation of ketones or phenols are desirable for the

development of technology for bio-oil upgrading.

Tetralin is an active hydrogen donor for the transfer hydrogenation of oxygenated compounds

(e.g., alkanals, ketones, and phenolics) on the low-cost acidic zeolite catalysts under mild

conditions (atmospheric pressure, 473-673 K), and tetralin itself is converted to naphthalene in

the reaction. On the other hand, tetralin can be produced by the naphthalene hydrogenation with

H2 at high pressures (>50 atm) on metallic catalysts (e.g., WS2+NiS+Al2O3 or

CoO+MoO3+Al2O3) [17]. Therefore, it is proposed that tetralin or alkyl tetralins can be applied

as hydrogen donor vehicles, which are recycled between the oxygenated compounds and H2,

allowing for the effective hydrodeoxygenation of these oxygenated species with an external

hydrogen source (H2) while avoiding the poisoning of metallic catalysts (for activating H2) with

the complex oxygenated species in bio-oil. The introduction of an external hydrogen source (H2)

for the hydrodeoxygenation of the oxygenated species would increase the effective H/C ratio,

and thus, would improve the heating value of the products [18]. In order to develop such a bio-

oil upgrading strategy, understanding the reaction chemistry, especially the transfer

190

hydrogenation, of additional oxygenated species within bio-oil (e.g., carboxylic acids, esters) on

solid Brønsted acid catalysts is necessary. In addition, efficient methods for separating the

hydrogen donor vehicles (e.g., naphthalene and tetralin) from the deoxygenation products are

required.

6.3. References



[2] C.B. Warren, A.B. Marin, J.F. Butler, Use of Unsaturated Aldehyde and Alkanol Derivatives

for Their Mosquito Repellency Properties, US5665781 A (1997).

[3] W. Pinkenhagen, α,β-Unsaturated Aldehydes and Their Use as Flavor-Modifying Ingredients,

US4381410 A (1983).

[4] W. Pickenhagen, A. Velluz, Flavoring with α, β-Unsaturated Aldehydes, US4324809 A

(1982).

[5] R. Veltri, G.B. Fodor, Pharmaceutically Useful Michael Addition Products of Unsaturated

Aldehydes and Ketones and Ascorbic Acid, US5098933 A (1992).


J. Catal., 341 (2016) 136-148.

[7] F. Lin, Y.-H. Chin, Catalytic Pathways and Kinetic Requirements for Alkanal Deoxygenation

on Solid Tungstosilicic Acid Clusters, ACS Catal., 6 (2016) 6634-6650.

[8] H. Wu, Q. Gong, D.H. Olson, J. Li, Commensurate Adsorption of Hydrocarbons and

Alcohols in Microporous Metal Organic Frameworks, Chem. Rev., 112 (2012) 836-868.



(2011) 257-268.







(1986).

191

[14] J.-S. Kim, Production, separation and applications of phenolic-rich bio-oil – A review,

Bioresour. Technol. , 178 (2015) 90-98.



389 (2010) 140-146.

[16] Y. Xue, S. Zhou, X. Bai, Role of Hydrogen Transfer during Catalytic Copyrolysis of Lignin

and Tetralin over HZSM-5 and HY Zeolite Catalysts, ACS Sustainable Chemistry &

Engineering, 4 (2016) 4237-4250.

[17] A.A. Krichko, D.V. Skvortsov, T.A. Titova, B.S. Filippov, N.E. Dogadkina, Production of

tetralin by the hydrogenation of naphthalene-containing fractions, Chem. Technol. Fuels Oils

5(1969) 18-22.

[18] N.Y. Chen, J.T.F. Degnan, L.R. Koenig, Liquid fuel from carbohydrates, Chem. Tech., 16

(1986) 506-511.

192

Appendix: Error Assessment

A1. Error assessment for rate measurements and

chemical titrations under identical conditions

The amounts of acid sites on the catalysts (e.g., MFI and FAU zeolites and

H4SiW12O40/SiO2 catalysts) were quantified using pyridine chemical titration. The

titration experiments were repeated for at least two times in order to assess the

measurement accuracy. For example, Table A1.1 summarizes the results of several

repeating pyridine titration experiments used for quantifying the amount of the initial

acid sites on the fresh H-MFI catalysts and the amount of remaining acid sites after

propanal reactions under various conditions, as reported in Chapter 2. The acid site

densities reported in this work are given by the average values (�̅�) measured from

these repeating measurements. The measurement errors are determined by the 95 %

confidence intervals:

Error = ±SE×1.96 (A1.1)

where 1.96 is the 0.975 quantile of the normal distribution; SE is the standard error of

the mean given by:

SE= s/√𝑛 (A1.2)

where s is the standard deviation and n is the number of measurements. The relative

errors are given by

Relative error= Error/�̅� (A1.3)

As shown in Table A1.1, the relative errors of these measurements are less than

±10 %. I also performed similar error assessments when measuring the acid site

densities on fresh H-FAU and H4SiW12O40/SiO2 catalysts, as well as the amount of

193

the remaining acid sites on these catalysts after alkanal reactions. The relative errors

are also within ±10 %.

Table A1.1. The results and error assessments for the measured acid sites densities on

fresh and spent H-MFI zeolites (Si/Al=11.5).

Reaction Condition

Acid site density (μmol gcat.‒1

) Relative

error

(%)

#1 #2 #3 Average Standard

deviation

Error

Fresh catalyst 1134 1086 1139 1120 24 ±27 ±2.4

Propanal reactiona 36.0 41.1 43.2 40.1 3.0 ±3.4 ±8.5

Propanal+H2O reaction 1b 59.6 59.9 59.7 0.1 ±0.2 ±0.3

Propanal+H2O reaction 2c 59.2 58.2 58.7 0.5 ±0.7 ±1.2

a 473K, 1.1 kPa propanal feed, space velocity=1.1 mmol (mol H+i·s)−1, time on stream=125 min;

b 473K, 1.1 kPa propanal+5 kPa H2O feed, space velocity=1.1 mmol (mol H+i·s)−1, time on stream=125 min;

c 1.1 kPa propanal+10 kPa H2O feed, 473 K, space velocity=1.1 mmol (mol H+i·s)−1, time on stream=125 min;

The reproducibility of the kinetic results reported in this work was assessed by

repeating the rate measurements under identical conditions (temperature, reactant

pressure, space velocity, and time-on-steam). For, example, Figure A1.1 shows the

measured rates for the pathways of inter- and intramolecular C=C bond formation

(rInter and rIntra, respectively) during propanal reaction on H-MFI zeolite at 473 K as a

function of time-on-stream. The measured rInter and rIntra values changed significantly

during the first 60 min of reactions. However, after the 125 min, the reactions were

approaching steady-state because the catalysts were deactivated to similar extents and

the propanal pressure became more identical (close to the feeding pressure) under low

conversion (<5 %), thus the deviation of the rate measurements reduced dramatically.

Therefore, I reported the reaction rates measured at steady-state (e.g., time on

stream=125 min) and used these values in the kinetic analysis. Table A1.2 lists the

measured rInter and rIntra at 125 min (given by the average values of the data in Figs.

194

A1.1a and A1.1b) and corresponding measurement errors (calculated via Eqns. A1.1-

A1.3). The relative errors are less than ±10 %, within the acceptable range.

Figure A1.1. Rates for (a) intermolecular C=C bond formation (rInter) and (b)

intramolecular C=C bond formation (rIntra) during propanal (1.1 kPa) reactions on H-

MFI zeolite (Si/Al=11.5) at 473 K as a function of time-on-stream (space

velocity=1.1×10−3

mol propanal (mol H+

i s)−1

).

Table A1.2. The results and error assessments for the measured rates of inter- and

intramolecular C=C bond formation (rInter and rIntra, respectively) during propanal

reaction on H-MFI zeolite at 473 K at 125 min.

Average Standard deviation Error Relative error (%)

rInter (10−8

mol (gcat.s)-1

) 0.814 0.073 ±0.071 ±8.7

rIntra (10−8

mol (gcat.s)-1

) 0.161 0.014 ±0.014 ±8.7

0 50 100 150 200 250 3000.0

0.5

1.0

1.5

2.0

r Inte

r (1

0-8

mo

l (g

ca

t. s

)-1)

Time on stream (min)

Run #1

Run #2

Run #3

Run #4

0 50 100 150 200 250 3000.0

0.5

1.0

1.5

r Intr

a (

10

-8 m

ol (g

ca

t. s

)-1)

Time on stream (min)

Run #1

Run #2

Run #3

Run #4

(a) (b)

195

A2. Error assessment for rate assessments at varying

alkanal pressures

Sometimes the rate measurements were carried out with different reactant pressures,

and these results were used for the linear or non-linear regression analysis. In this case,

the measurement errors were assessed based on the residual sums of squares (SR) of

the corresponding regression analysis. For example, for the kinetic data in Figures

2.5-2.7 in Chapter 2, I obtained the residual sums of squares (SR) for the respective

linear or non-linear regression using OriginPro 8.5. The standard deviations (s) of the

data were estimated by the square roots of the residual mean squares (s2), calculated

via Equation A1.4:

2 RSs s

n m

(A1.4)

where n‒m is the residual degree of freedom. The measurement errors were calculated

by Eqns. A1.1-A1.3, and listed in Table A1.3. Most of the standard deviations are less

than ±10 % of the measured values.

Table A1.3. Error assessments for the measured rates in Figures 2.5-2.7 of Chapter 2

(the rates are given in 10-5

mol (H+

r s)-1

).

Figures Measurement

range

Degree of

freedom

Residual sum

of squares

Standard

deviation

Error

Fig. 2.5 rinter 2.09-7.70 3 0.373 0.35 ±0.35

rintra 0.502-0.575 3 0.0298 0.031 ±0.031

Fig. 2.6 rinter 1.48-2.12 4 0.38 0.10 ±0.08

rintra 0.390-0.449 5 0.030 0.024 ±0.019

Fig. 2.7 rinter 0.585-2.28 5 0.0383 0.088 ±0.065

rintra 0.23-0.45 5 0.00228 0.02 ±0.02

196

A3. Error assessment for the measured rates of alkanal

transfer hydrogenation

The rate of alkanal (CnH2nO) transfer hydrogenation by a specific H-donor (RDH2)

(denoted as D2 2TH,C H O-R Hn n

r ) was measured indirectly. It was given by the difference

between the rates of intramolecular C=C bond formation in CnH2nO-RDH2 feed

mixture (D2 2Intra,C H O-R Hn n

r ) and CnH2nO feed (2Intra,C H On n

r ):

D D2 2 2 2 2TH,C H O-R H Intra,C H O-R H Intra,C H On n nn n nr r r (A1.5)

The error for 2Intra,C H On n

r was calculated (via Eqns. A1.1-A1.3) based on the standard

deviation of the repeated measurements under an identical condition as shown in

Figure A1.2. Table A1.3 summarizes the errors of 2Intra,C H On n

r at different time-on-

streams. These errors decrease as the time-on-stream increases and all the relative

errors are less than ±11 %.

D2 2Intra,C H O-R Hn nr were measured under different H-donor pressures (

D 2R HP ). It is

plausible to assume that D2 2Intra,C H O-R Hn n

r has a similar error level as 2Intra,C H On n

r .

According to Equation A1.5, the calculated transfer hydrogenation rate

D2 2TH,C H O-R Hn nr would have a propagated error (

D2 2TH,C H O-R Hn nError ) of:

2 2

D D2 2 2 2 2TH,C H O-R H Intra,C H O-R H Intra,C H On n nn n nError Error Error (A1.6)

where D2 2Intra,C H O-R Hn n

Error and 2Intra,C H On n

Error represent the errors of

D2 2Intra,C H O-R Hn nr and

2Intra,C H On nr , respectively.

197

Because all the D2 2Intra,C H O-R Hn n

r and 2Intra,C H On n

r values used in calculating

D2 2TH,C H O-R Hn nr were measured after 125 min, their measurement errors were

estimated to be less than ±0.017 μmol (gcat. s)-1

, as shown in Table A.1.4. Assuming

their errors were ±0.015μmol (gcat. s)-1

, the error for D2 2TH,C H O-R Hn n

r was estimated to

be ±0.021μmol (gcat. s)-1

, according to Equation A1.6.

Table A1.5 summarizes the measurement range for the rates of butanal transfer

hydrogenation with various H-donors (RDH2) D2 2TH,C H O-R Hn n

r on H-FAU at 573 K, as

reported in Chapters 3 and 4. When using a highly active H-donor (e.g., RDH2=tetralin

or cyclohexadiene), the measurement error of D2 2TH,C H O-R Hn n

r (±0.021 μmol (gcat. s)-1

)

is one order of magnitude lower than the D2 2TH,C H O-R Hn n

r values (0.13~0.25 μmol (gcat.

s)-1

). However, for a weak H-donor (RDH2=cyclohexene), the error of D2 2TH,C H O-R Hn n

r

is of the same magnitude of the D2 2TH,C H O-R Hn n

r values (0.020~0.022 μmol (gcat. s)-1

),

indicating a much lower accuracy for the kinetic measurement using a weak H-donor.

0 100 200 300 4000.0

0.1

0.2

0.3

0.4

0.5

0.6 Run #1

Run #2

Run #3

Run #4

Run #5

Run #6

r Intr

a,C

4H

8O

(m

ol (g

cat. s

)-1)


198

Figure A1.2. Rates for intramolecular C=C bond formation during butanal reaction

on H-FAU zeolite (4 8Intra,C H Or ) at 573 K as a function of time-on-stream (1.1 kPa

butanal, space velocity=7.4×10−3

mol butanal (mol H+

s)−1

).

Table A1.4. Error assessments for the measured rates for intramolecular C=C bond

formation during butanal reaction on H-FAU zeolites (Si/Al=15) at 573 K at different

time-on-streams.

Time-on-stream (min) 95 125 275 305 335

rIntra,C4H8O average (μmol (gcat. s)-1

) 0.227 0.195 0.115 0.114 0.111

Standard deviation 0.023 0.019 0.011 0.011 0.003

Error ±0.018 ±0.017 ±0.013 ±0.011 ±0.003

Relative error (%) ±8 ±9 ±11 ±10 ±3

Table A1.5. Measurement range for the rates of butanal transfer hydrogenation with

various H-donors (RDH2) (D2 2TH,C H O-R Hn n

r ) on H-FAU zeolites (Si/Al=15) at 573 K.

RDH2 rTH,C4H8O-RDH2 measurement range (μmol (gcat. s)-1

)

Cyclohexadiene 0.17~0.25

Tetralin 0.13~0.21

Cyclohexene 0.020~0.022

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