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Carbon-hydrogen bond activation: radical methane functionalization; unactivated alkene coupling; saccharide degradation; and carbon dioxide hydrogenation
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Carbon-hydrogen bond activation: radical methane functionalization; unactivated alkene coupling; saccharide degradation; and carbon dioxide hydrogenation
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Content
CARBON-HYDROGEN BOND ACTIVATION:
RADICAL METHANE FUNCTIONALIZATION; UNACTIVATED ALKENE
COUPLING; SACCHARIDE DEGRADATION; AND CARBON DIOXIDE
HYDROGENATION
by
Nima Zargari
A Dissertation Presented to the
FACULTY OF THE USC GRADUATE SCHOOL
UNIVERSITY OF SOUTHERN CALIFORNIA
In Partial Fulfillment of the
Requirements for the Degree
DOCTOR OF PHILOSOPHY
(CHEMISTRY)
May 2016
Copyright 2016 Nima Zargari
ii
ד״סב
Dedicated to
Maman Mahin
iii
ACKNOWLEDGEMENTS
It is due to the guidance of many integral people in my life that I have reached
this milestone. I would like to thank Dr. Donald R. Blake for the countless hours spent
with me during office hours, Dr. Gregory A. Weiss for sparking my interest in organic
chemistry, Dr. Andrej Lupták for opening the door to my first research opportunity, and
Dr. Kyung Woon Jung, for not giving up on me and continuously pushing me to
excellence.
I am also indebted to my close friends and family for their continuous support and
encouragement. I would like to thank my lab mate, Richard A. Giles, for his insightful
discussion. My closest friends, Paul Lapis; Daniel Lavi; Joshua Berookhim; and Michael
Broukhim, whom I consider to be my brothers. My grandfather, Farajollah Beroukhim,
who taught me patience. My sister and brother-in-law, Neda and Alfred Lavi, for their
unwavering support. My niece and nephews, Emunah, Emet, and Ehad Lavi, for
reminding me what is truly important in life. And my father, Rouhollah Zargari, for his
confidence in me.
I would also like to thank my mother, Forough Beroukhim. I am in great debt to
her undying support, boundless encouragement, and unconditional love. Thank you.
iv
Drop by drop, the river is formed.
v
TABLE OF CONTENTS
Dedication ii
Acknowledgments iii
Epigraph iv
Table of Contents v
List of Figures vii
List of Schemes ix
List of Tables x
Abstract xi
Chapter 1: A Mechanistic Insight into a Latent Radical Methane Process Propagated
by Trifluoromethyl Radicals 1
1.1 Introduction and Background 1
1.2 Results and Discussion 2
1.2.1 Discovery of Background Process 2
1.2.2
13
C Labeled Methane wet1D
1
H and
13
C NMR Experiments 4
1.2.3 Effects of TFA/TFAA 6
1.2.4 Formation Routes of Trifluoromethyl Radical 7
1.2.5 Proposed Methane Functionalization Propagated by
Trifluoromethyl Radicals 10
1.2.6 Density Functional Theory Calculations 12
1.2.7 Effects of Radical Inhibitors and Initiators 15
1.2.8 Effects of Varying Amounts of Hydrogen Peroxide 18
1.2.9 Summary 19
1.3 Experimental Methods 19
1.4 Chapter 1 References 21
Chapter 2: Hydroalkenylation: Palladium Catalyzed Coupling Reaction of
Unactivated Alkenes 24
2.1 Introduction and Background 24
2.2 Results and Discussion 25
2.2.1 Screening of Varying Palladium and Silver Sources 25
2.2.2 Screening of Varying Tetrafluoroborate and Boron Trifluoride
Additives 26
vi
2.2.3 Catalytic Complexation Between Palladium and BF
3
27
2.2.4 Effects of Tetrafluroborate and Boron Trifluoride Loading 28
2.2.5 Solvent Effect of AgBF
4
and HBF
4
29
2.2.6 Screening of Styrene Derivatives 30
2.2.7 Styrene-d
8
Study 31
2.2.8 Proposed Mechanism via a Palladium Hydride Complex 31
2.2.9 Summary 32
2.3 Experimental Methods 33
2.4 Chapter 2 References 34
Chapter 3: Conversion of Saccharides into Formic Acid using Hydrogen Peroxide
and a Recyclable Palladium(II) Catalyst in Aqueous Alkaline Media at Ambient
Temperature 36
3.1 Introduction and Background 36
3.2 Results and Discussion 37
3.2.1 Introduction to NHC-amidate Pd(II) Complex 37
3.2.2 Effects of Varying Amounts of Base and Time 39
3.2.3 Effects of Varying Amounts of Hydrogen Peroxide 40
3.2.4 Effects of Palladium Catalysts 40
3.2.5 Reaction Time Study 42
3.2.6 Effects of Catalyst Recycling 43
3.2.7 Mono- and Disaccharide Reducing Substrate Scope 43
3.2.8 Non-Reducing Saccharide Substrate Scope 44
3.2.9 Proposed Pd(II) assisted α-oxidative degradation pathway 45
3.2.10 Summary 47
3.3 Experimental Methods 47
3.4 Chapter 3 References 49
Chapter 4: Carbon Dioxide Hydrogenation: Catalysis via NHC-Amidate Metal
Complexes 51
4.1 Introduction and Background 51
4.2 Results and Discussion 53
4.2.1 Effects of Varying Temperature and Time 53
4.2.2 Effects of Varying Base 55
4.2.3 Effects of Palladium Salts 56
4.2.4 Effects of Metals Salts in with NHC-amidate Ligand 57
4.2.5 Effects at Pressure and Heterogeneous Catalyst Conditions 58
4.2.6 Proposed Mechanism via a Palladium-Hydroxide Complex 59
4.2.7 Summary 61
4.3 Experimental Methods 61
4.4 Chapter 4 References 62
vii
Bibliography 64
Appendices 73
Appendix 1: Supporting Information for Chapter 1 73
Appendix 2: Supporting Information for Chapter 2 104
Appendix 3: Supporting Information for Chapter 3 136
Appendix 4: Supporting Information for Chapter 4 155
viii
LIST OF FIGURES
Figure 1.1 wet1D
1
H NMR of
13
C labeled methane reaction 5
Figure 1.2
13
C NMR of
13
C labeled methane reaction 6
Figure 1.3
19
F NMR of of three methods of forming trifluoromethyl radical sources 10
Figure 1.4 Energy diagram from DFT calculations of Scheme 1.6 reactions 14
Figure 1.5 Transition state structure of reaction between trifluoromethyl radical
(·CF
3
) and methane 15
Figure 1.6 Transition state structure of reaction between the methyl radical and TFAA 16
Figure 1.7 Transition state structure of decomposition reaction of 5 into
trifluoromethyl radical and 6 17
Figure 2.1 Structure of NHC-amidate Pd(II) complex 24
Figure 3.1 Structure of NHC-amidate Pd(II) complex 38
Figure 3.2 Reaction time study 42
Figure 3.3 Effects of catalyst recycling 43
Figure 4.1 Structure of NHC-amidate Pd(II) complex 52
Figure 4.2 Time study preformed under standard reaction conditions from 1-6 hours 55
ix
LIST OF SCHEMES
Scheme 1.1 Selective examples of methane oxidation by transition metal catalysts
using hydrogen peroxide in TFAA 2
Scheme 1.2 Preliminary results on methane oxidation 3
Scheme 1.3 Background radical reaction at low temperature 3
Scheme 1.4 Formation of trifluoroperacetic acid and subsequent reaction with methane 8
Scheme 1.5 Formation of bis(perfluoroacetyl) peroxide and subsequent hemolysis 9
Scheme 1.6 Methane functionalization by proposed radical processes and termination
of trifluoromethyl radical 11
Scheme 2.1 Initial reaction conditions utilizing catalyst 1 25
Scheme 2.2 Catalytic complexation between palladium and BF
3
28
Scheme 2.3 Styrene-d
8
study 31
Scheme 2.4 Proposed mechanism via a palladium hydride complex 32
Scheme 3.1 Initial reaction conditions utilizing catalyst 1 38
Scheme 3.2 Proposed Pd(II) assisted α-oxidative degradation pathway of D-glucose
with hydrogen peroxide 46
Scheme 4.1 Initial reaction conditions utilizing catalyst 1 53
Scheme 4.2 Proposed mechanism through a palladium-hydroxide complex 60
x
LIST OF TABLES
Tables 1.1 Effects of TFA and TFAA 7
Table 1.2 Effects of radical inhibitors and initiators 18
Table 1.3 Effect of hydrogen peroxide 18
Table 2.1 Screening of varying palladium and silver sources 26
Table 2.2 Screening of varying tetrafluoroborate and boron trifluoride additives 27
Table 2.3 Effects of tetrafluoroborate and boron trifluoride loading 29
Table 2.4 Solvent effect with AgBF
4
and HBF
4
29
Table 2.5 Screening of styrene derivatives with HBF
4
30
Table 3.1 Effects of varying amounts of base and time 39
Table 3.2 Effects of varying amount of hydrogen peroxide 40
Table 3.3 Effects of palladium catalysts and loading 42
Table 3.4 Mono- and disaccharide reducing substrate scope 44
Table 3.5 Scope of reaction conditions with non-reducing substrates 45
Table 4.1 Effects of varying temperature and time 54
Table 4.2 Effects of varying base 56
Table 4.3 Effects of palladium catalysts 57
Table 4.4 Effects of varying metal salts and versatility of the NHC-amidate ligand 58
Table 4.5 Effect of varying palladium sources and solvent amount 59
xi
ABSTRACT
Due to the composition of organic compounds, the activation of carbon and
hydrogen bonds has grasped the attention of many modern chemists. The focus of the
following projects was centered around the activation and better understanding of C-H
bonds. From mechanistic studies of greenhouse gasses, such as methane and carbon
dioxide, co-dimerization of small molecules, or degradation of abundant biomass, the
breakage of C-H bonds and subsequent reformation of new C-H bonds are of paramount
importance.
In Chapter 1, thorough mechanistic studies and density functional theory (DFT)
calculations revealed a background radical pathway latent in metal catalyzed oxidations
of methane. Use of hydrogen peroxide with trifluoroacetic anhydride (TFAA) generated a
trifluoromethyl radical (·CF
3
), which in turn reacted with methane gas selectively to yield
acetic acid. It was found that the methyl carbon of the product was derived from methane
and the carbonyl carbon was derived from TFAA. Computational studies also support
these findings, revealing each step of the cyclic reaction is energetically favorable.
A highly efficient co-dimerization of styrene and cyclopentene, in the presence of
palladium and a BF
3
source, was developed selectively forming a new C-C bond, in
Chapter 2. The complex [Pd(PPh
3
)
2
]
+
BF
4
-
is believed to generate palladium hydride (Pd-
H), which catalyzes the reaction between various styrenes and cyclopentene in excellent
yields as single isomers. This co-catalytic system provides a new, efficient C-C bond
forming method.
xii
The development of an effective method that converts a variety of mono- and
disaccharides into predominantly formic acid is discussed in Chapter 3. A recyclable
NHC-amidate palladium(II) catalyst facilitates oxidative degradation of carbohydrates
without using excess oxidant. Stoichiometric amounts of hydrogen peroxide and sodium
hydroxide were employed at ambient temperatures.
Lastly, Chapter 4 discusses a highly efficient hydrogenation of carbon dioxide
into formic acid in the presence of the same NHC-amidate palladium(II) complex.
Excellent turnover number is observed when the catalyst was used under heterolytic
conditions. This catalytic system provides a new efficient carbon dioxide hydrogenation
method.
1
CHAPTER 1
A Mechanistic Insight into a Latent Radical Methane Process
Propagated by Trifluoromethyl Radicals
1.1 Introduction and Background
Methane, in the form of natural gas, is the most abundant hydrocarbon but the
least reactive due to the high bond-dissociation energy.
[1]
Because methane and its flared
product, carbon dioxide, are major greenhouse gases,
[2]
it is of paramount importance to
discover novel ways to functionalize methane other than the existing syngas generation
protocols. Currently, the most predominant methods to functionalize methane employ
sulfur-containing oxidants such as potassium persulfate
[3]
and sulfuric acid.
[4]
The most
ideal oxidant, molecular oxygen, has also been used to a lesser extent for the oxidation of
lower alkanes including methane.
[5]
In addition, various oxidants with transition metals or
even main-group catalysts with varying degrees of success have been investigated.
[6]
Hydrogen peroxide has been touted by many researchers in the oxidative
functionalization of methane because it is highly reactive and provides environmentally
friendly by-products. As illustrated in Scheme 1.1, Sen, Mizuno, and Ingrosso used
hydrogen peroxide in trifluoroacetic acid (TFA) and trifluoroacetic anhydride (TFAA)
successfully to convert methane gas to either methyl trifluoroacetate or acetic acid as
major liquid products.
[7-9]
Despite low yields based on methane, these methodologies
were impressive and pioneering. These transformations required the use of transition
metal catalysts; however, their mechanisms including the possible presence of non-
2
metallic background processes were not comprehensively studied. Reported is a
background reaction, which has been identified as a radical process.
[10]
The purpose of
this investigation is to provide clarity to the mysterious and underdeveloped reaction of
methane gas under radical conditions in TFAA/TFA.
1.2 Results and Discussion
1.2.1 Discovery of Background Process
Using various palladium catalysts including our NHC-amidate complex,
[11]
we
examined the oxidation of methane under conditions similar to the work done by Sen,
and obtained methanol in similar yields as Sen reported (Scheme 1.2). However, we
observed a mixture of products such as formic acid as the major product, carbon dioxide,
and acetic acid in small amounts. These reactions would occur via electrophilic
substitution as Sen pointed out, while the detection of multiple products, particularly
CH
4
(68 atm)
TFAA, H
2
O
2
CH
4
(30 atm)
CH
4
(50 atm)
Sen
Ingrosso
Mizuno
Pd(hfacac)
2
TFAA, H
2
O
2
Pd(CF
3
CO
2
)
2
80 °C, 24 h
TFAA, H
2
O
2
H
4
PV
1
Mo
11
O
40
80 °C, 24 h
75 °C, 4 h
CH
3
OTFA
CH
3
CO
2
H
CH
3
OTFA
(12%)
(3.2%)
(14%)
CH
3
CO
2
H
(3.5%)
Scheme 1.1 Selective examples of methane oxidation by transition metal catalysts using
hydrogen peroxide in TFAA
3
acetic acid, implied the presence of additional mechanistic pathways or background
reactions. To probe the presence of any background reactions, we carried out various
control experiments and made efforts to identify intermediates and products. When
methane was subjected to the oxidation by hydrogen peroxide in the absence of a
palladium source, both methanol (methyl trifluoroacetate before work-up) and acetic acid
were produced in 0.06% and 1.2% yields based on methane, respectively. Elevated
amounts of carbon dioxide and reduced amounts of acetic acid were observed at reactions
performed at 90 °C.
As depicted in Scheme 1.3, acetic acid was selectively generated from methane at
60 °C while another observed product, fluoroform stemmed from TFAA/TFA. Moreover,
less carbon dioxide was observed at low temperatures, and acetic acid was furnished as
the major or exclusive liquid product from methane at low temperatures such as 60 °C in
higher yields than at high temperatures (i.e., 90 - 100 °C).
Pd catalyst
TFAA, TFA, H
2
O
2
90 °C, 16 h
CH
4
w/o Pd catalyst
TFAA, TFA, H
2
O
2
90 °C, 16 h
CH
4
CH
3
OH CH
3
CO
2
H +
(major)
(1.2%) (0.06%)
CH
3
OH + CH
3
CO
2
H + HCO
2
H
Scheme 1.2 Preliminary results on methane oxidation
TFAA, TFA, H
2
O
2
60 °C, 16 h
CH
3
CO
2
H CF
3
H + CH
4
(3.8%)
Scheme 1.3 Background radical reaction at low temperature
4
1.2.2
13
C Labeled Methane wet1D
1
H and
13
C NMR Experiments
In the absence of methane, no products were observed, indicating that methane
was the actual carbon source for the intended reactions. When we employed
13
C labeled
methane, we observed acetic acid as the only meaningful
13
C incorporating product
(Figure 1.1). However, the acetyl protons in the
1
H NMR spectra exhibited a large one-
bond
13
C-H coupling (J
C-H
= 130 Hz) but no two-bond coupling (i.e.,
13
C-
13
C-H),
signaling only the methyl of the acetyl group came from methane. In addition, fluoroform
was constantly detected in variable amounts, suggesting that the C-C bond of the
trifluoroacetyl group was cleaved to offer the carbonyl moiety, which would be
incorporated into the produced acetic acid (vide infra).
5
Figure 1.1 wet1D
1
H NMR of
13
C labeled methane reaction
In the proton-coupled
13
C NMR spectra (Figure 1.2), similar patterns were
demonstrated, where the methyl carbon of the acetyl group showed a quartet with a large
coupling constant due to the aforementioned one-bond coupling (J
C-H
= 130 Hz). The
carbonyl carbon peak was negligible due to low abundance of
13
C, suggesting this carbon
would be naturally abundant
12
C. Both spectra confirm that methane was a key reactant to
form the methyl part and the carbonyl moiety would stem from another carbon source.
6
Figure 1.2
13
C NMR of
13
C labeled methane reaction
1.2.3 Effects of TFA/TFAA
The effect of TFAA and TFA were investigated to determine which reagent
participated in the formation of acetic acid (Table 1.1). The reaction did not proceed
when TFAA was excluded (entry 1). As the amount of TFAA increased, the reaction
afforded higher yields of acetic acid (entries 2 – 4). Use of pure TFAA resulted in a lower
yield suggesting acidic anhydrous conditions are required (entry 5). These results
supported the role of TFAA rather than TFA as the carbonyl source of the observed
acetic acid, especially since acetic acid was not observed in the absence of TFAA.
7
Table 1.1 Effects of TFA and TFAA
Entry TFAA / TFA (mmol) AcOH (µmol, yield) MeOH (µmol)
1 0.0 / 11 - -
2
3
4
5
1.4 / 8.1
2.8 / 5.4
4.3 / 2.7
5.7 / 0.0
33 (1.0%)
55 (1.7%)
121 (3.8%)
25 (0.77%)
2.0
0.7
-
-
Reaction conditions: Varying amounts of TFAA and TFA were added to 130 µmol 30%
H
2
O
2
in a stainless steel reactor with a high pressure valve and charged with 3200 µmol
methane. The mixture was stirred at 60 °C for 16 hours. D
2
O was added to the reaction
mixture and a wet1D NMR was taken using DMSO as an internal standard. Yields were
calculated based on methane as the limiting reagent.
1.2.4 Formation Routes of Trifluoromethyl Radical
Since CF
3
H was observed in both
1
H and
19
F NMR, we proposed that
trifluoromethyl radical (·CF
3
) are the chain carriers of this reaction. In conditions of
hydrogen peroxide and TFAA, there are several ways in which ·CF
3
can be produced. As
depicted in Scheme 1.4, the radical processes can be initiated by trifluoroperacetic acid 1
(TFPAA), which is in accordance with previous studies performed with TFAA and
hydrogen peroxide at room temperature.
[12]
Therefore, TFPAA (1) can react directly with
methane to form intermediate 2.
[13]
Radical degradation of this complex would furnish
methyl (·CH
3
) and TFA radicals (3) along with water. Decarboxylation of TFA radical 3
leads to the facile production of trifluoromethyl radical (·CF
3
), which plays a crucial role
in the subsequent radical propagation processes. Interestingly, previous studies have
shown that trifluoroacetyl peroxides furnish fluoroform along with carbon dioxide when
they were subject to thermal decomposition in hydrocarbon solvents.
[14]
8
Another source of ·CF
3
is bis(perfluoroacetyl) peroxide. When hydrogen peroxide
is added to a solution of TFAA and trace amounts of base, 1 is no longer produced and
the primary product is bis(perfluoroacetyl) peroxide 4 (Scheme 1.5).
[12]
Alternatively,
when 30% hydrogen peroxide is introduced dropwise into a biphasic system of an
aqueous alkaline media and Freon 113, 4 is produced and the organic layer which may be
separated from the aqueous solution.
[15]
At higher temperatures, 4 readily undergoes
homolysis, and the resulting TFA radical 3 rapidly dissociates to give rise to the
formation of trifluoromethyl radical (·CF
3
).
Our standard reaction conditions where applied using all three different pathways
of producing trifluoromethyl radical initiators and all produced acetic acid as the selective
product. Most importantly, in the reaction which bis(perfluoroacetyl) peroxide in Freon
113 solution was used as the trifluoromethyl radical initiator, no additional hydrogen
peroxide was introduced into the solution. This supports our belief that the radical chain
carrier of this methane functionalization reaction is trifluoromethyl radical (·CF
3
).
Scheme 1.4 Formation of trifluoroperacetic acid and subsequent reaction with methane
TFAA
30% H
2
O
2
25 °C
H
O
O
O
CF
3
2
CF
3
+ CH
3
- H
2
O - CO
2
H
O
O
O
CF
3
H
H
3
C
1
CH
4
3
F
3
C O
O
9
As shown in Figure 1.3, the
19
F NMR of the solution prior to reacting them with
methane clearly show the formation of either 1 or 4. Since our standard reaction
conditions do not use base, then trifluoroperacetic acid is the trifluoromethyl radical
source of the reactions.
Scheme 1.5 Formation of bis(perfluoroacetyl) peroxide and subsequent homolysis
TFAA
30% H
2
O
2
, 0.1% KOH
25 °C
O O CF
3
F
3
C
O O
bis(trifluoroacetyl) peroxide
TFAA
Na
2
CO
3
, NaCl, H
2
O
30% H
2
O
2
, Freon 113, -3 °C
O O CF
3
F
3
C
O O
bis(trifluoroacetyl) peroxide
in Freon 113 solution
3
4
F
3
C O
O
O CF
3
O
CF
3
- CO
2
10
Figure 1.3
19
F NMR of three methods of forming trifluoromethyl radical sources
1.2.5 Proposed Methane Functionalization Propagated by
Trifluoromethyl Radicals
As illustrated in Scheme 1.5, trifluoromethyl radical (·CF
3
) can react with
methane to form methyl radical (·CH
3
) and fluoroform (CF
3
H) because of the favorable
difference in the bond dissociation energies between fluroform (446.4 kJ/mol) and
methane (439.7 kJ/mol).
[16]
Theoretical calculations show that methyl radical (·CH
3
) are
10.3 kJ/mol more stable than trifluoromethyl radical (·CF
3
).
[17]
After the produced methyl
radical (·CH
3
) adds to TFAA, the resultant intermediate 5 can undergo radical
11
dissociation to mixed anhydride 6, while trifluoromethyl radical (·CF
3
) can be
regenerated to propagate the radical cycle. Upon treatment with water during the work-
up, mixed anhydride 6 is ultimately liberated into acetic acid and TFA. This mechanistic
scheme is similar to the one proposed by Sen while investigating the radical-initiated
oxidative functionalization of higher chain alkanes such as ethane and propane.
[7b]
Termination of this radical process may occur by the reaction of two trifluoromethyl
radicals (·CF
3
), forming hexafluoroethane (C
2
F
6
). Hexafluoroethane is commonly seen as
a decomposition product of reactions containing trifluoromethyl radicals (·CF
3
).
[18]
This mechanistic approach accounts for a few critical observations. First, greater
amounts of acetic acid were observed as the amount of TFAA increased because TFAA
provided both CF
3
radical as the reagent and the carbonyl group as the substrate.
Secondly, acetic acid was not detected until the aqueous work-up because the mixed
anhydride 6 was stable during the reaction. As a consequence, the over-oxidation of
CF
3
O F
3
C
O O
CF
3
O F
3
C
O O
CH
3
CF
3
CH
3
O F
3
C
O O
CF
3
H
5
6
CH
4
CH
3
CF
3
+ CF
3
C
2
F
6
Termination
H
2
O
AcOH
+ TFA
Scheme 1.6 Methane functionalization by proposed radical processes and termination of
trifluoromethyl radical
12
acetic acid to carbon dioxide would be inhibited. Lastly, since water could be formed as a
by-product of the radical degradation of complex 2, excess amounts of TFAA were used
to maintain anhydrous reaction conditions.
Seeking evidence to corroborate the proposed mechanism, we carefully analyzed
products by using NMR and GC methods mainly to identify important product species
such as fluoroform, hexafluoroethane, carbon dioxide, and mixed anhydride 6. We
confirmed the presence of fluoroform by both wet1D
1
H (Figure 1.1) and
19
F NMR.
Hexafluoroethane was observed as another fluorine containing product besides
fluoroform by
19
F NMR analysis. By employing wet1D
1
H and
13
C NMR techniques, we
detected the mixed anhydride at the end of the reaction as well as its rapid hydrolysis to
acetic acid and TFA upon the addition of water. Under standard reaction conditions
(Table 1, entry 4), 109 µmol of carbon dioxide was produced, which was verified by gas
chromatography analysis in addition to proton-coupled
13
C NMR studies (Figure 1.2).
These results suggested that 130 µmol of hydrogen peroxide led to 109 µmol of CO
2
and
the equimolar amount of trifluoromethyl radical (·CF
3
). Acetic acid as the only liquid
product was obtained in slightly larger amounts (i.e., 121 µmol), implying the radical
cycle proposed in Scheme 5 would be in play.
1.2.6 Density Functional Theory Calculations
Rigorous computational studies were also performed to support these findings.
The radical processes for methane functionalization in Scheme 1.6, with the exception of
the hydrolysis of mixed anhydride 6 and the radical termination reaction, were analyzed
13
computationally using Gaussian 09.
[13]
The bimolecular reaction between the
trifluoromethyl radical (·CF
3
) and methane will be referred to as reaction 1. The
bimolecular addition reaction between the methyl radical (·CH
3
) and TFAA will be
referred to as reaction 2. Finally, the unimolecular decomposition reaction of 5 into
trifluoromethyl radical and 6 will be referred to as reaction 3. Each of these reactions
proceeds through a single transition state which will be referred to as transition state 1, 2,
and 3 respectively. In all cases the reactant, product, and transition state geometries were
optimized using the B3LYP hybrid functional,
[16]
and Dunning’s aug-cc-pVDZ basis
set.
[17]
Diffuse functions were included in the basis set in order to account for long-range
interactions, particularly between the fluorine and carbon atoms.
Energies were calculated for the gas-phase optimized geometries both with and
without the conductor-like polarizable continuum solvent model (COSMO).
[19-21]
The
COSMO model was used to estimate the solvent effects of TFA on the single-point
energies of all reactants, transition states, and products. The COSMO results predict that
the energies and free energies of reaction for reactions 1 and 3 are lowered by 5.0 kJ/mol
and 10.1 kJ/mol, respectively, when using TFA as a solvent but increased by 1.8 kJ/mol
for reaction 2. For all reactions, the energies of activation did not change significantly
when using the COSMO model. From these data we conclude that solvation plays an
important role in the reaction thermodynamics of Scheme 1.5, if not the kinetics (Figure
1.3).
14
Figure 1.4 Energy diagram from DFT calculations of Scheme 1.6 reactions
Frequency calculations were performed on the optimized geometries to obtain
zero-point and thermal corrections at 333K and 27.0 atm, and to confirm that our
transition state geometries are first-order saddle points and that our reactant and product
geometries are minima on the potential energy surface. The single imaginary frequency
calculated for each transition state was verified to correspond to the reaction coordinate
for that step.
0
62
-8.1
58.9
2.5
40.9
-52.6
Relative Energy (kJ/mol)
CF
3
+
CH
4
H CH
3
F
3
C
CF
3
H +
CH
3
CF
3
O F
3
C
O O
CH
3
CF
3
O F
3
C
O O
CH
3
5
CF
3
O F
3
C
O O
CH
3
CH
3
O F
3
C
O O
6
15
The transfer of a hydrogen atom to proceed from ·CF
3
and CH
4
to CF
3
H and ·CH
3
is found to occur through transition state 1 (Figure 1.4). The C-H bond length in this
structure is 1.37 Å between ·CF
3
and the shared hydrogen atom and 1.33 Å between ·CH
3
and the shared hydrogen atom. The FCH bond angle is approximately 109.8° and the
HCH bond angle is approximately 104.3°. The solvated energy of activation for this
transition state was found to be 39.3 kJ/mol and one negative frequency was calculated at
-1604 cm
-1
with an IR intensity of 300 km/mol.
Figure 1.5 Transition state structure of reaction between trifluoromethyl radical (·CF
3
)
and methane
Of the three reactions modeled, the transition state for reaction 2 has the lowest
solvated activation energy of 22.6 kJ/mol, indicating that this is the fastest reaction in the
mechanism. The predicted structure of transition state 2 shows that addition of the methyl
radical ·CH
3
has a low energy barrier because the electronegative fluorine atoms on the
TFAA create a strong electrophilic carbon center (Figure 1.5). This is consistent with
16
the decrease in spin density on the ·CH
3
carbon and an increase in spin density on the
carbonyl oxygen of TFAA. The C-C bond distance in this structure is 2.86 Å between the
TFAA carbon atom and the ·CH
3
carbon atom. The CCO bond angle between the
previously mentioned carbon atoms and the respective carbonyl oxygen atom is
approximately 100.5°. Frequency analysis on this transition state structure found one
negative frequency at -319 cm
-1
with an IR intensity of 67 km/mol.
Figure 1.6 Transition state structure of reaction between the methyl radical and TFAA
The structure of transition state 3 shows the removal of the ·CF
3
group from the
reactant to form the mixed anhydride product (Figure 1.6). The C-C bond length in this
structure is 2.07 Å between ·CF
3
and the carbonyl carbon of the mixed anhydride. The
FCC bond angle between the ·CF
3
moiety and the carbonyl carbon is approximately
106.0°. The solvated energy of activation for this transition state was found to be 40.2
kJ/mol and one negative frequency was calculated at -358 cm
-1
with an IR intensity of 6
km/mol.
17
Figure 1.7 Transition state structure of decomposition reaction of 5 into trifluoromethyl
radical and 6
1.2.7 Effects of Radical Inhibitors and Initiators
In an effort to validate the radical process, we evaluated the feasibility of the
methane functionalization by using various radical inhibitors and promoters (Table 1.2).
When various radical inhibitors and initiators including butylated hydroxytoluene (BHT),
TEMPO, and azobisisobutyronitrile (AIBN) were added (10 mole percent to hydrogen
peroxide), yields of acetic acid were reduced by 75% (entries 2 – 4 vs entry 1). As
expected, a stoichiometric amount of BHT shut down the reaction completely (entry 5).
When radical initiators such as AIBN were solely utilized in the absence of hydrogen
peroxide, no products were observed. These data would support a radical mechanism and
the necessity of TFPAA (1).
18
Table 1.2 Effects of radical inhibitors and initiators
Entry H
2
O
2
(µmol) Additive
AcOH (µmol) MeOH
(µmol)
1 130 - 121 -
2
3
4
5
130
130
130
130
BHT
a
TEMPO
a
AIBN
a
BHT
b
34
22
46
-
1.4
2.1
4.3
-
a
13 µmoles.
b
130 µmoles. Reaction conditions: 2.7 mmol TFA, 4.25 mmol TFAA, 130
µmol 30% H
2
O
2
, and varying amounts of an radical inhibitor and initiator were added to
a stainless steel reactor with a high pressure valve and charged with 3200 µmol methane.
The mixture was stirred at 60 °C for 16 hours. D
2
O was added to the reaction mixture and
a wet1D NMR was taken using DMSO as an internal standard.
1.2.8 Effects of Varying Amounts of Hydrogen Peroxide
As represented in Table 1.3, we varied the amounts of hydrogen peroxide and
compared the aforementioned results (shown again in Table 1.3, entry 2). Reducing the
amount of hydrogen peroxide by half produced acetic acid in a lower yield as expected
(entry 1). However, doubling the amount of hydrogen peroxide failed to furnish
additional acetic acid (entry 3). Both of these modified amounts of hydrogen peroxide
produced methanol in trace amounts. Therefore, there seemed to be an appropriate
stoichiometry between hydrogen peroxide and TFAA for optimal radical processes (entry
2).
Table 1.3 Effect of hydrogen peroxide
Entry H
2
O
2
(µmol) AcOH (µmol, yield) MeOH (µmol)
1 65 75 (2.4%) 1.9
2
3
130
260
121 (3.8%)
109 (3.4%)
-
1.8
Reaction conditions: 2.7 mmol TFA, 4.25 mmol TFAA were added to varying amounts
of 30% H
2
O
2
and were added to a stainless steel reactor with a high pressure valve and
charged with 3200 µmol methane. The mixture was stirred at 60 °C for 16 hours. D
2
O
was added to the reaction mixture and a wet1D NMR was taken using DMSO as an
internal standard.
19
1.2.9 Summary
A background radical reaction latent in metal catalyzed oxidative
functionalization of methane has been founded and extensively studied. Hydrogen
peroxide together with TFAA produced the trifluoromethyl radical (·CF
3
), which in turn
activated the methane C-H bond. The resulting methyl radical reacted with TFAA to
generate a mixed anhydride, which was subsequently hydrolyzed into acetic acid. NMR
data indicated that the methyl group of acetic acid originated from methane while the
carbonyl carbon originated from TFAA. Experimental results from the use of radical
initiators and inhibitors along with careful analysis of products and intermediates
supported our hypothesis on a radical mechanism. Though similar reaction conditions for
methane oxidation have frequently been employed, a thorough study of this intriguing
background reaction has not been performed. Thus, we hope that this report will bring
some clarity to radical transformations and metal catalyzed oxidations of methane, the
most abundant hydrocarbon.
1.3 Experimental Methods
To a 0.5 dram vial equipped with a stir bar was added a solution of 0.2 mL (2.69
mmol) TFA, 0.6 mL (4.25 mmol) TFAA, and 13 µL (130 µmol) 30% H
2
O
2
. The solution
was stirred at room temperature for one minute. The vial was then placed in a stainless
steel reactor with a high-pressure valve composed of Swagelok components and was
purged and charged with 27 atm (3200 µmol) methane gas. The mixture was stirred at
20
60 °C for 16 hours. 0.5 mL D
2
O was added to the reaction mixture and a wet1D NMR
was taken using DMSO as an internal standard.
21
1.4 Chapter 1 References
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Han, X.; Bao, X. J. Am. Chem. Soc. 2006, 128, 16028-16029. (j) Zerella, M.;
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Eur. J. 2015, 21, 1286-1293. (b) Guo, Z.; Liu, B.; Zhang Q.; Deng, W.; Wang, Y.; Yang,
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Bischof, S. M.; Gustafson, S. J.; Devarajan, D.; Gunsalus, N.; Ess, D. H.; Periana, R. A.
Science 2014, 343, 1232-1237. (d) Gupta, S.; Kirillova, M. V.; da Silva, M. F. C. G.;
Pombeiro, A. J. L.; Kiri, A. M. Inorg. Chem. 2013, 52, 8601-8611. (e) Silva, T. F. S.;
Mac Leod, T. C. O.; Martinsa, L. M. D. R. S.; da Silva, M. F. C. G.; Schiavon, M. A.;
Pombeiro, A. J. L. J. Mol. Catal. A: Chem. 2013, 367, 52-60. (f) Hammond, C.; Forde,
M. M.; Ab Rahim, M. H.; Thetford, A.; He, Q.; Jenkins, R. L.; Dimitratos, N.; Lopez-
Sanchez, J. A.; Dummer, N. F.; Murphy, D. M.; Carley, A. F.; Taylor, S. H.; Willock,
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D. J.; Stangland, E. E.; Kang, J.; Hagen, H.; Kiely, C. J.; Hutchings, G. J. Angew. Chem.
Int. Ed. 2012, 51, 5129-5133. (g) Reis, P. M.; Silva, J. A. L.; Palavra, A. F.; Fraústo da
Silva, J. J. R.; Kitamura, T.; Fujiwara, Y.; Pombeiro, A. J. L. Angew. Chem. Int. Ed.
2003, 42, 821-823. (h) Lin, M.; Hogan, T.; Sen, A. J. Am. Chem. Soc. 1997, 119, 6048-
6053. (i) Lin, M.; Hogan, T.; Sen, A. J. Am. Chem. Soc. 1996, 118, 4574-4580. (j)
Nakata, K.; Yamaoka, Y.; Miyata, T.; Taniguchi, Y.; Takaki, K.; Fujiwara, Y. J.
Organomet. Chem. 1994, 473, 329-334. (k) Vargaftik, M. N.; Stolarov, I. P.; Moiseev, I.
I. J. Chem. Soc., Chem. Commun. 1990, 1049-1050.
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Hogan, T.; Sen, A. J. Am. Chem. Soc. 1997, 119, 2642-2646.
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5944.
[9] Ingrosso, G.; Midollini, N. J. Mol. Catal. A: Chem. 2003, 204-205, 425-431.
[10] Radical processes were reported using persulfate and peroxodisulfate: (a) Basickes,
N.; Hogan, T. E.; Sen, A. J. Am. Chem. Soc. 1996, 118, 13111-13112. (b) Lin, M.; Sen,
A. J. Chem. Soc., Chem. Commun. 1992, 892-893.
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Chem. Int. Ed. 2008, 47, 9326-9329.
[12] Krasutsky, P. A.; Kolomitsyn, I. V.; Carlson, R. M. Org. Lett. 2001, 3, 2997-2999.
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Chem. 2001, 66, 789-795.
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(Engl. Transl.) 1980, 16, 1360-1362.
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23
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24
CHAPTER 2
Hydroalkenylation: Palladium Catalyzed Coupling Reaction of
Unactivated Alkene
2.1 Introduction and Background
Seminally, the discovery of palladium catalyzed cross-coupling reactions have
forever changed the field of modern organic chemistry.
[1]
Reactions of palladium
catalysts with tetrafluoroborate additives have been used to activate and functionalize
various alkenes although these are underexplored.
[2]
Sen,
[3]
Shmidt,
[4]
and Tkach,
[5]
amongst other, have utilized this co-catalytic system in order to dimerize or polymerize
various olefins; however, they have not been able to couple two different alkenes
efficiently. Herein we report a vinyl arene coupling reaction utilizing a palladium(0)
source and a tetrafluoroborate co-catalyst. In sharp contrast to previous work, this
reaction is a chemoselective and stereoselective cross-coupling of different alkenes
without the need of a leaving or activating group on the olefins.
Figure 2.1 Structure of NHC-amidate Pd(II) complex
N
N
O
N
O
Me
Cl
Pd
Me
1
25
Initially, we used our NHC-amidate Pd(II) complex 1 (Figure 2.1) by activating
with AgBF
4
to remove the chlorine ligand and form a cationic palladium(II) complex.
[6]
Under mild conditions, we were able to successfully couple styrene and cyclopentene in a
high yield, while minimizing the homo-coupled side product of styrene (Scheme 2.1).
2.2 Results and Discussion
2.2.1 Screening of Varying Palladium and Silver Sources
We then endeavored to determine the most favorable source of palladium by
screening different palladium and silver complexes (Table 2.1). Similarly to 1, other
palladium sources provided the desired product as well. Palladium(II) sources such as
Pd(OAc)
2
and Pd(TFA)
2
(entries 1 and 2), yielded low to good results, respectively.
Pd(dba)
2
(entry 3) was slightly less effective as Pd(TFA)
2
; however, Pd(PPh
3
)
4
(entry 4)
produced the desired product in comparable yields as 1. No reaction took place in the
absence of a palladium source (entry 5) or in the absence of AgBF
4
(entry 6). Using
AgPF
6
as the silver source drastically lowered the yield of the reaction (entry 7), while
other silver(I) salts provided no catalytic activity (entries 8-9). This led us to discover the
significance of AgBF
4
.
Ph
+
1 (5 mol%), AgBF
4
(10 mol%)
1,2-DCE, 50 °C, 16 hrs.
100 µmol
50 eq.
Ph
95%
Scheme 2.1 Initial reaction conditions utilizing complex 1
26
Table 2.1 Screening of varying palladium and silver sources
a)
Entry Palladium Source Silver Source Yield (%)
b)
1 Pd(OAc)
2
AgBF
4
37
2
3
Pd(TFA)
2
Pd(dba)
2
AgBF
4
AgBF
4
79
73
4 Pd(PPh
3
)
4
AgBF
4
93
5 - AgBF
4
-
6 Pd(PPh
3
)
4
- -
7 Pd(PPh
3
)
4
AgPF
6
13
8 Pd(PPh
3
)
4
AgOTf -
9 Pd(PPh
3
)
4
AgNO
3
-
a)
Reaction conditions: 100 µmol styrene, 1,250 µmol cyclopentene, 5 µmol palladium
source, and 20 µmol silver source were dissolved in 0.7 mL 1,2-DCE and stirred at 50 °C
for 16 hours.
b)
Determined by
1
H NMR using DMSO as an internal standard.
2.2.2 Screening of Varying Tetrafluoroborate and Boron Trifluoride
Additives
Since AgBF
4
was previously used to remove chlorine in 1, its necessity with
Pd(PPh
3
)
4
was tested by screening Pd(PPh
3
)
4
against varying tetrafluoroborate and boron
trifluoride additives (Table 2.2). NaBF
4
by itself or in the presence of acids such as
hydrochloric or p-toluenesulfonic acid provided no products (entries 1-3). However,
NaBF
4
in the presence of acids such as methanesulfonic or sulfuric acid provided the
product in low yields (entries 4-5). Importantly, in the absence of the tetrafluoroborate
salt, no products were observed when methanesulfonic or sulfuric acid were used alone
(entries 6-7). The best results were observed when HBF
4
was used as the
tetrafluoroborate additive, producing the product in 85% yield (entry 8). Further, boron
trifluoride diethyl etherate was also used as an trifluoride additive, which produced a
27
significant amount of the desired product, revealing that the reaction may not solely be
catalyzed by fluoroboric acid (entry 9). Critically, employing other Lewis acids such as
AlCl
3
did not produce any product (entry 10). These results strongly suggested that the
tetrafluoroborate or boron trifluoride additives were responsible for activating the
catalytic system.
Table 2.2 Screening of varying tetrafluoroborate and boron trifluoride additives
a)
Entry Additive 1 Additive 2 Yield (%)
b)
1 NaBF
4
- -
2
3
NaBF
4
NaBF
4
HCl
TsOH
-
-
4 NaBF
4
MsOH 15
5 NaBF
4
H
2
SO
4
35
6 - MsOH -
7 - H
2
SO
4
-
8 HBF
4
- 85
9 BF
3
·OEt
2
- 76
10 AlCl
3
- -
a)
Reaction conditions: 100 µmol styrene, 1,250 µmol cyclopentene, 5 µmol Pd(PPh
3
)
4
, 20
µmol additive 1, and 20 µmol additive 2 were dissolved in 0.7 mL 1,2-DCE and stirred at
50 °C for 16 hours.
b)
Determined by
1
H NMR using DMSO as an internal standard.
2.2.3 Catalytic Complexation Between Palladium and BF
3
Based on these results and the work done by Shmidt, et. al., the trace amounts of
fluorine in the solution form palladium-fluorine dimers (Scheme 2). With increasing
amounts of BF
3
, the active catalytic species is formed, wherein BF
3
is complexed to the
palladium source either via a fluorine atom (F·BF
3
) or as a BF
4
-
anion.
[7]
28
2.2.4 Effects of Tetrafluroborate and Boron Trifluoride Loading
The ratio between palladium source and the boronic cocatalyst have been
extensively studied in the dimerization or oligomerization of alkenes. The loading of both
tetrafluoroboric acid and boron trifluoride were examined under our standard reaction
conditions (Table 3). Stoichiometric equivalents of the boronic source to the palladium
catalyst produced low yields (entry 1). Doubling the ratio of tetrafluoroboric acid
significantly increased the yield, while doubling the ratio of boron trifluoride modestly
increased the yield of the desired product (entry 2). A four-time excess of either
tetrafluoroborate and boron trifluoride loading to the palladium catalyst produced the
highest amount of the desired product (entry 3), while the yield of the reaction decreased
with higher ratio amounts (entry 4). These results further support our belief that the
activated catalyst is a palladium source complex with a BF
4
-
anion, since increased
amounts of the boronic cocatalyst would shift the equilibrium of the palladium complex
towards more of an ionic character species.
Pd
F
F
Pd
+ BF
3
- BF
3
Pd
F
L
BF
3
+ L
- L
Pd
L
L
2
BF
4
-
2
Scheme 2.2 Catalytic complexation between palladium and BF
3
29
Table 2.3 Effects of tetrafluoroborate and boron trifluoride loading
a)
Entry HBF
4
(µmol) Yield (%)
b)
BF
3
⋅OEt
2
(µmol) Yield (%)
b)
1 5 24 5 8
2
3
10
20
66
85
10
20
20
76
4 40 49 40 23
a)
Reaction conditions: 100 µmol styrene, 1,250 µmol cyclopentene, 5 µmol Pd(PPh
3
)
4
,
and varying amounts of HBF
4
or BF
3
·OEt
2
were dissolved in 0.7 mL 1,2-DCE and stirred
at 50 °C for 16 hours.
b)
Determined by
1
H NMR using DMSO as an internal standard.
2.2.5 Solvent Effect of AgBF
4
and HBF
4
The solvent effect of both AgBF
4
and HBF
4
were then tested separately (Table
2.4). Both AgBF
4
and HBF
4
reacted similarly in polar and nonpolar halogenated solvents
(entries 1-3). In the case of using benzene as the solvent, HBF
4
produced moderate
yields, while the same reaction using AgBF
4
did not produce the desired product (entry
4). Other polar solvents inhibited the reaction (entries 5-8). This may be due to the
solvent’s interaction between BF
3
, deactivating the cocatalyst.
Table 2.4 Solvent effect with AgBF
4
and HBF
4
a)
Entry Solvent AgBF
4
Yield (%)
b)
HBF
4
Yield (%)
b)
1 DCM
c)
86 88
2
3
1,2-DCE
chloroform
93
90
85
74
4 benzene - 59
5 acetonitrile - -
6 THF - -
7 MeOH - -
8 1,2-dioxane - -
a)
Reaction conditions: 100 µmol styrene, 1,250 µmol cyclopentene, 5 µmol Pd(PPh
3
)
4
,
and 20 µmol AgBF
4
or HBF
4
were dissolved in 0.7 mL solvent and stirred at 50 °C for 16
hours.
b)
Determined by
1
H NMR using DMSO as an internal standard.
c)
Reaction run at
25 °C.
30
2.2.6 Screening of Styrene Derivatives
We then examined the electronics of the reaction by varying the substituents on
styrene in order to determine how electron withdrawing or donating groups would affect
the yield of the desired product. These reactions were tested with HBF
4
as the boronic
cocatalyst (Table 5). Interestingly, electron deficient styrene substrates produced the
highest amount of the desired product (entries 2-6). Electron rich styrene substrates
produced drastically lower yields (entries 7-8), while the most electron rich substituent,
4-methoxystyrene, produced the lowest yield. Regardless of acid loading and temperature
variations, 4-methoxystyrene polymerized under these reaction conditions. Therefore,
electron deficient styrene substrates yielded higher results in these reaction conditions not
due to their inherent electron characteristics, but since electron rich styrene substrates are
too reactive and produce polymerized side products instead.
Table 2.5 Screening of styrene derivatives with HBF
4
a)
Entry Substrate Yield (%)
b)
1 styrene 85
2
3
4-trifluoromethylstyrene
4-nitrostyrene
96
91
4 4-chlorostyrene 91
5 2-fluorostyrene 82
6 4-fluorostyrene 80
7 4-methylstyrene 40
8 2-vinylnaphthalene 57
9 4-methoxystyrene 14
a)
Reaction conditions: 100 µmol styrene substrate, 1,250 µmol cyclopentene, 5 µmol
Pd(PPh
3
)
4
, and 20 µmol HBF
4
were dissolved in 0.7 mL 1,2-DCE and stirred at 50 °C for
16 hours.
b)
Determined by
1
H NMR using DMSO as an internal standard
31
2.2.7 Styrene-d
8
Study
Preforming the reaction using deuterated styrene revealed interesting mechanistic
implications (Scheme 3). Critically, the final product was not deuterated on the
cyclopentene ring. The deutrerated ratio at the benzylic position of the product was 2:1,
revealing that there were no proton or deuterium shifts during the reaction.
2.2.8 Proposed Mechanism via a Palladium Hydride Complex
Sen et al. suggested a carbocationic mechanism for the dimerization of styrene,
using Pd(PPh)
2
(BF
4
)
2
in the late 1980s;
[3b]
however, this theory has been long challenged.
Such a carbocationic mechanism involves the participation and abstraction of free H
+
and
would be very sensitive to the temperature fluctuations. Later findings have refuted this
belief and proposed that the coupling of unactivated alkenes with similar catalytic
systems may be propagated by palladium hydride complexes
[8]
or hydrido-palladium
species.
[9]
Though hydrido-palladium species selectively provide head-to-tail products,
such as our product, we favor the palladium hydride complex due to its simplicity and
long list of precedence.
It is well known that Brønsted acids, such as HBF
4
, complex with Pd(0)
+
Pd(PPh
3
)
4
, BF
3
·OEt
2
1,2-DCE, 50 °C, 16 hrs.
D
D
D
D
CD
2
H
D
D
H
Scheme 2.3 Styrene-d
8
study
32
sources, forming palladium hydride complexes, e.g. HPd(L
2
)BF
4
.
[10]
Scheme 4 proposes
the coupling of styrene and cyclopentene via a palladium hydride complex.
[8,11]
Initially,
the coordinately unsaturated palladium hydride species I coordinates to the alkene of
styrene, while two monodentate phosphine ligands also coordinate to the palladium
center, forming structure II. The palladium then forms a π complex with the abundant
cyclopentene, forming structure III and initiating the cross-coupling reaction. Lastly,
beta-elimination of IV terminates the chain, providing the desired product and
regenerating the palladium hydride active species I.
2.2.9 Summary
In conclusion, we have found a set of reaction conditions, which selectively
couple styrene and cyclopentene, providing a head-to-tail cross-coupling product. In the
L
2
Pd
Ph
PdL
2
Ph
Ph
Ph
H
PdL
2
L
2
Pd H
Ph
I
II
III
IV
Scheme 2.4 Proposed mechanism via a palladium hydride complex
33
presence of palladium and a BF
3
source, the active catalytic species is formed, wherein
BF
3
is complexed to the palladium source. Palladium hydride complexes initiate the
reaction and are more selective towards electron poor substrates.
2.3 Experimental Methods
To a 1-dram vial equipped with a stir bar, were added 100 µmol of substrate, 1,250 µmol
of cyclopentene, 5 µmol of Pd(PPh
3
)
4
, 20 µmol of AgBF
4
or HBF
4
diethyl ether complex
(51-57%). The mixture was dissolved in 0.7 mL 1,2-DCE and stirred at 50 °C for 16
hours. The solution was then passed through a silica column, filtered using hexanes, and
the solvent was evaporated under reduced pressure. The
1
H NMR of the remaining oil
was taken using CDCl
3
as the NMR solvent and DMSO as an internal standard.
34
2.4 Chapter 2 References
[1] a) Lyons, T. W.; Sanford, M. S. Chem. Rev. 2010, 110, 1147-1169; b) Chen, X.;
Engle, K. M.; Wang, D. H.; Yu, J. Q. Angew. Chem. 2009, 121, 5196-5217; Angew.
Chem. Int. Ed. 2009, 48, 5094-5115; c) Littke, A. F.; Fu, G. C. Angew. Chem. 2002, 114,
4350-4386; Angew. Chem. Int. Ed. 2002, 41, 4176-4211; d) Miyaura, N.; Suzuki, A.
Chem. Rev. 1995, 95, 2457-2483; e) Stille, J. K. Angew. Chem. Int. Ed. Engl. 1986, 25,
508-524.
[2] a) Oehme, G.; Pracejus, H. J. Organomet. Chem. 1987, 320, C56-C58; b) Oehme, G.;
Pracejus, H. Tetrahedron Lett. 1979, 4, 343-344; c) Kaneda, K.; Terasawa, M.; Imanaka,
T.; Teranishi, S. Tetrahedron Lett. 1977, 34, 2957-2958.
[3] a) Jiang, Z.; Sen, A. Organometallics 1993, 12, 1406-1415; b) Sec, A.; Lai, T. W.;
Thomas, R. R. J. Organomet. Chem. 1988, 358, 567-588; c) Sen, A.; Lai,T. W.
Organometallics. 1982, 1, 415-417.
[4] a) Tkach, V. S.; Myagmarsuren, G.; Suslov, D. S.; Darjaa, T.; Dorji, D.; Shmidt, F. K.
Catal. Commun. 2008, 9, 1501-1504; b) Tkach, V. S.; Myagmarsuren, G.; Suslov, D. S.;
Mesyef, M.; Shmidt, F. K. Catal. Commun. 2007, 8, 677-680; c) Myagmarsuren, G.;
Tkach, V. S.; Suslov, D. S.; Shmidt, F. K. Russ. J. Appl. Chem. 2007, 80, 252-256; d)
Myagmarsuren, G.; Tkach, V. S.; Suslov, D. S.; Chernyshev, M. L.; Shmidt, F. K. React.
Kinet Catal. Lett. 2007, 90, 137-143; e) Tkach, V. S.; Suslov, D. S.; Gomboogiin, M.;
Ratovskii, G. V.; Shmidt, F. K. Russ. J. Appl. Chem. 2006, 79, 85-88; f) Tkach, V. S.;
Zelinskii, S. N.; Ratovskii, G. V.; Proidakov, A. G.; Shmidt, F. K. Russ. J. Appl. Chem.
2004, 30, 703-708; g) Tkach, V. S.; Myagmarsuren, G.; Mesyef, M.; Shmidt, F. K. React.
Kinet. Catal. Lett. 1999, 66, 281-287; h) Chernyshev, M. L.; Tkach, V. S.; Dmitrieva, T.
V.; Ratovskii, G. V.; Zinchenko, S. V.; Shmidt, F. K. Kinet. Catal. 1997, 38, 527-531; i)
Chernyshev, M. L.; Tkach, V. S.; Dmitrieva, T. V.; Ratovskii, G. V.; Zinchenko, S. V.;
Shmidt, F. K. React. Kinet. Catal. Lett. 1992, 48, 291-294; j) Tkach, V. S.; Shmidt, F. K.;
Murasheva, N. A.; Malakhova, N. D.; Dmitrieva, T. V.; Ratovskii, G. V. React. Kinet.
Catal. Lett. 1988, 36, 213-216.
[5] a) Suslov, D. S.; Pahomova, M. V.; Abramov, P. A.; Bykov, M. V.; Tkach, V. S.
Catal. Commun. 2015, 67, 11-15; b) Suslov, D. S.; Bykov, M. V.; Belova, M. V.;
Abramov, P. A.; Tkach, V. S. J. Organomet. Chem. 2014, 752, 37-43; c) Kuratieva, N.
V.; Tkach, V. S.; Suslov, D. S.; Bykov, M. V.; Gromilov, S. A. J. Struct. Chem. 2011, 52,
813-815; d) Tkach, V. S.; Suslov, D. S.; Kuratieva, N. V.; Bykov, M. V.; Belova, M. V.
Russ. J. Coord. Chem. 2011, 37, 752-756; e) Suslov, D. S.; Tkach, V. S.; Bykov, M. V.;
Belova, M. V. Pet. Chem. 2011, 51, 157-163.
[6] Lee, J. H.; Yoo, K. S.; Park, C. P.; Olsen, J. M.; Sakaguchi, S.; Prakash, G. K. S.;
Mathew, T.; Jung, K. W. Adv. Synth. Catal. 2009, 351, 563-568.
35
[7] a) Tkach, V. S.; Suslov, D. S.; Myagmarsuren, G.; Ratovskii, G. V.; Rohin, A. V.;
Felix, T.; Shmidt, F. K. J. Organomet. Chem. 2008, 693, 2069-2073; b) Myagmarsuren,
G.; Tkach, V. S.; Suslov, D. S.; Shmidt, F. K. React. Kinet. Catal. Lett. 2005, 85, 197-
203; c) Myagmarsuren, G.; Tkach, V. S.; Shmidt, F. K. React. Kinet. Catal. Lett. 2004,
83, 337-343; d) Bobkova, A. V.; Zelinskii, S. N.; Ratovskii, G. V.; Tkach, V. S.; Shmidt,
F. K. Kinet. Catal. 2001, 42, 189-192; e) Tkach, V. S.; Shmidt, F. K.; Ratovskii, G. V.;
Malakhova, N. D.; Murasheva, N. A.; Chernyshev, M. L.; Burlakova, O. V. React. Kinet.
Catal. Lett. 1988, 36, 257-262.
[8] Myagmarsuren, G.; Tkach, V. S.; Shmidt, F. K.; Mohamad, M.; Suslov, D. S. J. Mol.
Catal. A: Chem. 2005, 235, 154-160.
[9] Ma, H.; Sun, Q.; Li, W.; Wang, J.; Zhang, Z.; Yang, Y.; Lei, Z. Tetrahedron Lett.
2011, 52, 1569-1573.
[10] a) Tkach, V. S.; Suslov, D. S.; Myagmarsuren, G.; Shmidt, F. K. Russ. J. Appl.
Chem. 2007, 80, 1419-1423; b) Leoni, P. J. Organomet. Chem. 1991, 418, 119-126; c)
Leoni, P.; Sommovigo, M.; Pasquali, M.; Midollini, S.; Braga, D.; Sabatino, P.
Organometallics. 1991, 10, 1038-1044.
[11] Choi, J. H.; Kwon, J. K.; RajanBabu, T. V.; Lim, H. J. Adv. Synth. Catal. 2013, 355,
3633-3638.
36
CHAPTER 3
Conversion of Saccharides into Formic Acid using Hydrogen Peroxide
and a Recyclable Palladium(II) Catalyst in Aqueous Alkaline Media at
Ambient Temperature
3.1 Introduction and Background
The conversion of biomass into valuable chemical feedstock products is an
emerging theme in the field of chemistry.
[1]
Notably, sugars comprise a majority of
biomass, of which glucose is the primary monosaccharide by mass.
[2]
Therefore, being
abundant carbon sources, it is of paramount importance to discover novel ways to
degrade such monosaccharides and convert them into valuable chemical feedstock
products such as 5-hydroxymethylfurfural (HMF), furfural, succinic acid, lactic acid, and
formic acid.
[3]
Numerous studies have shown that formic acid can be employed as a
hydrogen source in direct formic acid fuel cells (DFAFC). Moreover, due to their ease of
refueling, efficiency, and safety, DFAFC are considered to be an alternative to methanol
and hydrogen fuel cells. Having such characteristics allows DFAFC to potentially be
utilized in common consumer electronics in addition to being a power source to
automobiles.
[4]
Therefore it is of great value to be able to convert abundant and
inexpensive biomass into formic acid efficiently and under mild conditions.
Various work has been performed in degrading sugars into useful chemical
feedstock products. Such degradation techniques include the use of acid, high-
temperature liquid water, singlet oxygen, and other oxidants. The most prominent
protocols utilized acidic media and produced HMF in moderate yields, where large
37
amounts of mineral acids such as hydrochloric and sulfuric acids were required.
[5]
Further, though the use of high-temperature liquid water in degrading sugars excludes the
use of catalysts, high pressures and temperatures of up to 10 MPa and 600 °C were
required in addition to a 2-4 molar excess of oxidant.
[6]
While diradical oxygen
degradation of reducing sugars sensibly employed the use of an abundant and
inexpensive oxidant, the lack of product selectivity hindered the practicality of this
methodology.
[7]
Solid-supported palladium catalysts have also been utilized in the
oxidation of glucose into gluconic acid; however, these catalysts need to be activated
between 300-500 °C in a hydrogen or argon atmosphere.
[8]
Several studies have employed hydrogen peroxide in the presence of an alkali and
transition metals in the degradation of monosaccharaides and oligosaccharides.
[9]
In
particular, Isbell et. al. thoroughly examined such conditions in great detail.
[10]
However,
high yields were achieved using excess amounts of base and oxidant and/or only after
long reaction times (>300 h). Described herein is the development of efficient catalytic
methods for the oxidative degradation of common saccharides to formic acid.
3.2 Results and Discussion
3.2.1 Introduction to NHC-amidate Pd(II) Complex
These conditions employ a novel NHC-amidate Pd(II) complex
[11]
1 (Figure 3.1)
and stoichiometric amounts of hydrogen peroxide under aqueous alkaline conditions at
ambient temperatures. In sharp contrast to previous work, this simple methodology
38
gives rise to the highly efficient formation of formic acid as an exclusive or predominant
product under mild conditions using the minimal amount of oxidant.
Figure 3.1 Structure of NHC-amidate Pd(II) complex
Hydrogen peroxide was the oxidant used in this procedure since it is convenient,
accessible, and breaks down into environment-friendly side products unlike other
oxidants. In an effort to perform these reactions in the most sustainable manner, water
was also chosen as the sole solvent in the reaction mixture. Under these conditions, our
NHC-amidate palladium complex remains stable, yet is highly active.
Utilizing NHC-amidate Pd(II) complex 1 in our initial reaction conditions
(Scheme 3.1) provided inefficient yet promising results. Though over half of the starting
material was consumed over the reaction time, there was only a 12 percent carbon mass
balance to either formic acid or glycolic acid. We believed that the excess amount of
hydrogen peroxide caused the overoxidation of formic acid into carbon dioxide.
Scheme 3.1 Initial reaction conditions utilizing catalyst 1
H OH
O
100 µmol D-Glucose
(0.2 M aqueous solution)
1 mmol 30% H
2
O
2
5 mol % 1, 6 h., 40 °C
5.6 TON
N
N
O
N
O
Me
Cl
Pd
Me
1
39
3.2.2 Effects of Varying Amounts of Base and Time
Therefore, in order to prevent this thermodynamically favorable reaction from
taking place, base was added to the reaction mixture, consequently converting the formic
acid produced into a more stable formate salt and preventing overoxidation of the product
(Table 3.1).
[6b-d]
The addition of NaOH to the reaction mixture increased the turnover of
formic acid 15-fold (entry 3) compared to a reaction in the absence of any base (entry 1).
The addition of KOH was less effective (entry 2) and resulted in catalyst degradation.
Lower amounts of base afforded proportionally less product despite longer reaction times
(entry 4). It was found that overnight treatment with six equivalents of NaOH led to the
highest yield of the formate salt (entry 5). This is reasonable since each molecule of
glucose should theoretically produce six molecules of formic acid. Additional amounts of
base did not change the outcome significantly (entry 6). A reaction was conducted with
sodium chloride in the absence of base to determine any potential role of the alkali metal.
This reaction did not produce any formic acid, suggesting the alkali metal does not
catalyze the reaction.
Table 3.1 Effects of varying amounts of base and time
Entry Base Amount (µmol) Time (h) HCOOH (TON)
a
1
2
-
KOH
-
600
6
6
5.6
44.9
3 NaOH 600 6 85.2
4
5
6
NaOH
NaOH
NaOH
300
600
1,200
16
16
16
48.1
119.8
110.1
a)
Determined by wet1D NMR with a DMSO standard. Reaction conditions: 100 µmol
glucose, 5 µmol 1, and varying amounts of base were dissolved in 0.4 mL H
2
O. 1.0 mmol
H
2
O
2
was added and the mixture was stirred at 40 °C.
40
3.2.3 Effects of Varying Amounts of Hydrogen Peroxide
In light of these impressive results, we decided to optimize conditions, and
investigated the effect of hydrogen peroxide (Table 3.2). We observed a background
reaction in the absence of hydrogen peroxide, which produced minimal amounts of
formic acid (entry 1). Using less than stoichiometric amounts of the hydrogen peroxide
furnished only a fraction of the formic acid expected (entry 2). Therefore, adding a
stoichiometric amount of hydrogen peroxide was necessary to afford formic acid at over
100 TON (entry 3). Excess amounts of oxidant increased the yield marginally (entry 4),
suggesting that stoichiometric amounts of hydrogen peroxide (six equivalents to each
glucose; one equivalent to each carbon) would be optimal for practical use.
Table 3.2 Effects of varying amount of hydrogen peroxide
Entry H
2
O
2
(µmol) HCOOH (TON)
a
1 - 6.2
2 250 76.4
3
4
600
1,000
109.7
114.6
a)
Determined by wet1D NMR with a DMSO standard. Reaction conditions: 100 µmol
glucose, 5 µmol 1, and 600 µmol NaOH were dissolved in a 0.5 mL H
2
O/ H
2
O
2
solution
and stirred at 25 °C for 16 hours.
3.2.4 Effects of Palladium Catalysts
Comparing complex 1 with other palladium salts produced interesting results
(Table 3.3). While each reaction using palladium salts consumed all the starting material,
catalyst 1 produced the highest TON of formic acid. Pd(MeCN)
2
Cl
2
, which was the
41
palladium precursor of catalyst 1, generated a small amount of formic acid (entry 1)
compared to catalyst 1 (entry 5). Pd(OAc)
2
yielded twice the amount of formic acid as
Pd(MeCN)
2
Cl
2
, yet the product carbon mass balance remained less than stoichiometric
(entry 2). PdCl
2
was the least effective of the palladium sources, degrading all of the
glucose but only producing two TON of formic acid (entry 3). Interestingly, it was
concluded that these palladium salts catalyzed the overoxidation of the formate salt into
carbon dioxide, since the amount of formic acid observed was significantly higher in the
absence of the palladium salts (entry 4). This hypothesis was verified by stirring sodium
formate under the standard reaction conditions. A significant amount (up to 80%) of the
formate was decomposed to carbon dioxide in the presence of the palladium salts, while
catalyst 1 did not affect the amount of formate, thus offering high chemoselectivity. In
addition, decreased catalyst 1 loading continued to produce significant amounts of formic
acid (entries 5-6) but did not completely convert the starting material to product.
Most importantly, the greatest differentiating factor between our NHC-amidate
Pd(II) catalyst and the palladium salts tested was the fact that our catalyst could be
recycled and reused through the simple means of extraction, while maintaining high
reactivity (vide infra). In the absence of catalyst 1, various Lewis acids were tested under
standard reaction conditions, all of which produced lower TON than catalyst 1 (see
supporting information). These results suggest that our NHC-amidate Pd(II) complex has
distinguished reactivity compared to commercially available Lewis acids.
42
Table 3.3 Effects of palladium catalysts and loading
Entry Pd Catalyst mol % HCOOH (TON)
a
1 Pd(MeCN
2
)Cl
2
5 22.7
2 Pd(OAc)
2
5 40.0
3
4
5
6
PdCl
2
-
1
1
5
-
2.5
1.25
2.1
76.3
192.7
390.0
a)
Determined by wet1D NMR with a DMSO standard. Reaction conditions: 100 µmol
glucose, varying amounts of a palladium source, and 600 µmol NaOH were dissolved in a
0.5 mL H
2
O/ H
2
O
2
solution and stirred at 25 °C for 16 hours.
3.2.5 Reaction Time Study
With optimal conditions in hand, we examined the time dependence of this
transformation (Figure 3.2). A high initial rate of conversion was followed by a gradual
degradation of the remaining starting material. Within the first three hours of the reaction,
about half of the formic acid was produced, while it took an additional 13 hours for the
reaction to go to completion.
Figure 3.2 Reaction time study
0
20
40
60
80
100
120
0 2 4 6 8 10 12 14 16 18
Formic Acid (TON)
Time (hours)
43
3.2.6 Effects of Catalyst Recycling
We then decided to examine the feasibility of catalyst reclamation. We found that
due to the NHC-amidate ligand of catalyst 1, at the end of each reaction the catalyst can
be extracted using methylene chloride and reprocessed in a sequential reaction while still
maintaining its reactivity and selectivity. We were able to recover and utilize our catalyst
three times and still produce formic acid at over 100 TON (Figure 3.3). This is of
paramount value since it allows for the removal of the heavy metal from the aqueous
media. Thus this procedure has the promise to be utilized in water purification and
reclamation while simultaneously producing a chemical feedstock product, which can be
used as an alternative fuel source.
Figure 3.3 Effects of catalyst recycling
3.2.7 Mono- and Disaccharide Reducing Substrate Scope
The developed conditions were applied to various carbohydrates (Table 3.4).
40
60
80
100
120
0 1 2 3
Formic Acid (TON)
Times Recycled
44
Other monosaccharide reducing substrates such as D-galactose, D-ribose, and D-xylose
garnered high yields of formic acid selectively (entries 1-3), while D-fructose and D-
tagatose furnished formic acid as a major product and significant amounts of glycolic
acid as the minor product (entries 4-5). This was believed to be due to the difference
between the structure and functionality of aldoses and ketoses. For disaccharide reducing
substrates, such as D-maltose, D-lactose, and D-cellobiose, the reaction temperature was
increased to 60 °C in order to achieve comparable amounts of formic acid overnight
(entries 6-8). Alternatively, running the reactions at room temperature for longer periods
of time (48 h) produced similar results.
Table 3.4 Mono- and disaccharide reducing substrate scope
Entry Substrate HCOOH (TON)
Glycolic Acid (TON)
1 D-Galactose 111.2 -
2 D-Ribose 119.7 -
3
4
5
6
7
8
D-Xylose
D-Fructose
D-Tagatose
D-Maltose
a
D-Lactose
a
D-Cellobiose
a
116.9
83.1
89.3
86.0
98.7
97.4
-
14.5
16.3
-
-
-
a)
Reaction was heated to 60 °C. Reaction conditions: 100 µmol glucose, 5 µmol 1, and
600 µmol NaOH were dissolved in a 0.5 mL H
2
O/ H
2
O
2
solution and stirred at 25 °C for
16 hours.
3.2.8 Non-Reducing Saccharide Substrate Scope
As expected, the conversion yields of formic acid with non-reducing substrates
were poor (Table 3.5), due to their lack of hemiacetal and hemiketal groups. Sucrose, D-
melezitose, and D-raffinose produced formic acid in the range of 15 – 30 TON (entries
45
1-3) compared to over 100 TON for reducing substrates. Moreover, increasing the
amount of oxidant and reaction time did not render higher conversions, as it did with
reducing substrates. Applying our reaction conditions to a sugar alcohol such as glycerol
did not produce significant yields either (entry 4). The low yields were not believed to be
due to the overoxidation of these substrates to carbon dioxide. A large amount of starting
material remained at the end of the reaction, indicating that the degradation of these
substrates was less active.
Table 3.5 Scope of reaction conditions with non-reducing substrates
Entry Substrate HCOOH (TON)
1 Sucrose 27.3
2 D-Melezitose 15.9
3
4
D-Raffinose
Glycerol
28.5
39.1
Reaction conditions: 100 µmol glucose, 5 µmol 1, and 600 µmol NaOH were dissolved in
a 0.5 mL H
2
O/ H
2
O
2
solution and stirred at 60 °C for 16 hours.
3.2.9 Proposed Pd(II) assisted α-oxidative degradation pathway
Mechanistically, there are two primary routes of the oxidative degradation of
aldoses into formic acid (Scheme 3.2). As reported in previous work, initial oxidation of
the aldehyde is most likely, followed by breaking of the C1-C2 bond (α-scission).
[10c, 12]
This forms the first equivalent of formic acid and a successive aldehyde at C2. This
process is continued until the complete degradation of the hexose and the formation of six
molecules of formic acid is produced. Alternatively, rather than initially cleaving the C1-
C2 bond, a β-scission cleaves the C2-C3 bond, yielding an equivalent of oxalic acid,
which in turn degrades to CO
2
and formic acid, as well as a shorter-chain aldose. Since
46
ketoses are slower to oxidize than aldoses, harsher conditions and lower yields of formic
acid were observed.
Based on our results, we believe our procedure proceeds primarily through a α-
oxidation mechanism. Our primary source of evidence is the nearly complete conversion
of monosaccharides into formic acid. Since the carbon turnover of the monosaccharides
used was 93% or higher, a mechanism entailing the loss of half of the carbon mass by
decarboxylation of oxalic acid
[13]
would not be viable. In addition, under alkaline aqueous
conditions with catalyst 1 and stoichiometric amounts of hydrogen peroxide, oxalic acid
did not produce formic acid. Lastly, only trace amounts of oxalic acid were detected by
Scheme 3.2 Proposed Pd(II) assisted α-oxidative degradation pathway of D-glucose with
hydrogen peroxide
HCOOH
α-oxidation
β-oxidation
+
6 HCOOH
HOOC-COOH
3 HCOOH + 3 CO
2
+
Major
pathway
Minor
pathway
2 3 4
5
6
HO
OH
OH
OH
O
H
3
4
5
6
HO
OH
OH
OH
O
1
2 3 4
5
6
[Pd]
HO
OH
OH
OH
O
HO
H
47
13
C NMR under our standard conditions after four hours. As suggested in our previous
work
[11]
catalyst 1 acts as a bidentate system, simultaneously activating an aldehyde and
alcohol in close proximity of each other, forming a five-membered heterocycle glucose
adduct. This enhances the nucleophilic attack of peroxide, promoting the α-oxidation in a
systematic sequential order. While other palladium salts may activate saccharides in a
similar fashion, catalyst 1 avoids the overoxidation of formate into carbon dioxide,
setting it apart from other Pd(II) counterparts.
3.2.10 Summary
A set of efficient and environmentally friendly conditions that oxidatively
degraded saccharides using stoichiometric amounts of hydrogen peroxide in aqueous
alkaline media at ambient temperatures have been discovered. The NHC-amidate Pd(II)
complex 1 catalyzed the reactions, likely acting as a Lewis acid and activating the
aldehyde functional group in the aldoses. Contrary to other procedures, this method did
not use excess amounts of oxidant and did not require the input of heat for aldoses, while
the catalyst was recyclable and maintained its efficiency. This methodology can become
of great value since sugars comprise a majority of biomass. Thus, our developed
conditions contribute to converting this abundant carbon source into alternative fuels.
3.3 Experimental Methods
100 µmol of substrate, 5 µmol of 1, and 600 µmol NaOH were dissolved in 0.44
48
mL H
2
O. 60 µL 30% H
2
O
2
was added and the mixture stirred at 25 °C for 16 hours. 0.25
mL of D
2
O was then added to the reaction mixture with a sealed capillary DMSO
standard. The solution was then analyzed using wet1D NMR.
49
3.4 Chapter 3 References
[1] Gallezot, P. Chem. Soc. Rev. 2012, 41, 1538-1558.
[2] Liu, D.; Nimlos, M. R.; Johnson, D. K.; Himmel, M. E.; Qian, X. J. Phys. Chem. A.
2010, 114, 12936-12944.
[3] Tong, X.; Ma, Y.; Li, Y. Appl. Catal., A. 2010, 385, 1-13.
[4] (a) Rees, N. V.; Compton, R. G. J. Solid State Electrocem. 2011, 15, 2095-2100; (b)
Uhm, S.; Chung, S. T.; Lee, J. J. Power Sources 2008, 178, 34-43; (c) Rice, C.; Ha, S.;
Masel, R. I.; Waszczuk, P.; Wieckowski, A.; Barnard, T. J. Power Sources 2002, 111,
83-89.
[5] (a) Qi, L.; Mui, Y. F.; Lo, S. W.; Lui, M. Y.; Akien, G. R.; Horváth, I. T. ACS Catal.
2014, 4, 1470-1477; (b) Román-Leshkov, Y.; Davis, M. E. ACS Catal. 2011, 1, 1566-
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51
CHAPTER 4
Carbon Dioxide Hydrogenation: Catalysis via NHC-Amidate Metal
Complexes
4.1 Introduction and Background
Being an undesirable side product of the industrial revolution and the driving
force of the destruction of biomass, carbon dioxide is a greenhouse gas which is difficult
to activate and react due to its thermodynamic stability.
[1]
Moreover, since there are
exhaustible resources available, it is of paramount importance to utilize carbon dioxide as
a C
1
-building block not only to reduce the amount of greenhouse gasses, but also to
recycle carbon dioxide as a basic raw material into chemical feedstock products or
alternative fuel sources. Various work has been performed in order to functionalize
carbon dioxide by either direct hydrogenation or activation and reaction of the gas.
Employing Salen metal catalysts, carbon dioxide has been copolymerized with epoxides
producing polycarbonates,
[2]
while cyclic carbonate and polycarbonate formation was
achieved by group 3, 4, and 5 metal complexes.
[3]
The most sought out reaction of carbon dioxide is the direct hydrogenation of the
gas into valuable products such as methanol, formamides, and formic acid.
[4]
Using
supercritical carbon dioxide and a ruthenium catalyst, carbon dioxide has been selectively
converted into methanol.
[5]
Other reaction conditions employ carbon dioxide at much
lower pressures yet still produce methanol in the presence of copper, zinc, and aluminum
co-catalysts.
[6]
In the presence of ammonia and primary and secondary alkyl amines,
52
the direct carbonylation of carbon dioxide may take place in order to produce formamides
or methylamines.
[7]
Performing the reactions in supercritical carbon dioxide significantly
increased the yields of these products, compared to reactions using liquid solvents.
A plethora of studies have been found that selectively hydrogenate carbon dioxide
into formic acid under varying solvents, metal catalysts, and bases. This is of great value
since several studies have employed formic acid as a hydrogen source in direct formic
acid fuel cells (DFAFC). DFAFC are considered to be an alternative to methanol and
hydrogen fuel cells due to their ease of refueling, efficiency, and safety. This allows
DFAFC to potentially be utilized in common consumer electronics and automobiles.
[8]
Notably, recent studies have employed the use of NHC ligands in order to reduce
carbon dioxide into either methanol or formic acid.
[9]
Herein is report the development of
a set of reaction conditions which selectively hydrogenate carbon dioxide into formic
acid utilizing our NHC-amidate Pd(II) complex 1 (Figure 4.1).
[10]
Figure 4.1 Structure of NHC-amidate Pd(II) complex
N
N
O
N
O
Me
Cl
Pd
Me
1
53
4.2 Results and Discussion
4.2.1 Effects of Varying Temperature and Time
Our initial reactions conditions (Scheme 4.1) closely emulated those of Takezaki
et. al. where a palladium source was utilized in the presence of base, producing formate
salts as the major product.
[11]
When an equal pressure of carbon dioxide and hydrogen
gas were stirred at 160 °C for three hours in the presence of 1 in an alkaline aqueous
solution, 32.2 µmols of formic acid were produced as the major product, while methanol
and formaldehyde were produced in trace amounts.
In light of these promising results we decided to further optimize the reaction
conditions in order to obtain the highest turnover number (TON) of the products based on
1. Therefore, since the molar ratio between carbon dioxide and hydrogen is 3 : 1 in order
to produce one equivalent of methanol and water, we decide to change the ratio of the
carbon dioxide and hydrogen gasses to 3 : 1 to better suite the stoichiometry.
Interestingly, the reaction selectively produced formic acid as the sole product, in
addition to doubling the yield (Table 4.1, entry 1). Running the reaction at the same
temperature over a longer period of time also significantly increased the yield (Table 4.1,
entry 2); however, it was found that increasing the temperature to 200 °C for the same
150 psi H
2
, 1 mL H
2
O CO
2
150 psi
1.7 mmol KOH, 5 µmol 1, 3 hrs.
HCOOH
32.2 µmol
MeOH
0.5 µmol
HCHO
0.2 µmol
Scheme 4.1 Initial reaction conditions utilizing complex 1
54
three hours produced the greatest results (Table 4.1, entry 3). Lowering the reaction
temperature exponentially decreased the amount of product formed (Table 4.1, entry 4-5).
Table 4.1 Effect of varying temperature and time
Entry Temp. (°C) Time (hrs.) HCOOH (µmols)
a
TON
1 160 3 65.8 12.0
2 160 16 176.1 35.2
3 200 3 311.3 56.6
4 100 3 26.7 4.9
5 60 3 4.6 0.8
a)
Determined by wet1D NMR with a DMSO standard. Reaction conditions: 5 µmol 1,
and 1.7 mmol KOH were dissolved in a 1.0 mL H
2
O. The reaction mixture was then
placed in a stainless steel reactor with a high-pressure valve composed of Swagelok
components and was purged and charged with 6.8 atm (806 µmol) carbon dioxide gas,
followed by 20 atm (2,370 µmol) hydrogen gas.
The time dependence of this transformation was then examined (Figure 4.2). The
reaction reaches the highest turnover number after three hours, while half of the product
amount is produced after two hours. Having the reaction run for an additional hour
drastically decreased the yield, presumably to the thermodynamic stability of the product.
Between 4-6 hours the interconversion equilibrated to about 170 µmol of formic acid.
This may be due to the interconversion between formic acid into H
2
and CO
2
, which was
previously reported by other Pd-NHC compounds or metals.
[12]
It is important to note that
running the reaction for six consecutive hours or three hours followed by recharging the
system with additional H
2
and CO
2
and running the reaction again for three hours
produced similar results.
55
4.2.2 Effects of Varying Base
Varying the base added to the reaction mixture was then tested. Notably, in the
absence of any base, no reaction took place (Table 4.2, entry 1). KOH not only stood out
from other group one alkaline bases (Table 4.2, entry 2-6), but also from group two
alkaline bases and organic bases as well. This may be due to not only the basicity of the
solution, but the size of the counter ion as well. The major product of the reaction using
calcium hydroxide was methanol in low yield (Table 4.2, entry 7), while silver
trifluoroacetate provided formaldehyde as the major product (Table 4.2, entry 8). Amine
bases such as triethylamine and ammonium hydroxide failed to produce any significant
result or produced trace amounts of both formic acid or methanol, respectively (Table
4.2, entry 9-10).
0
50
100
150
200
250
300
0 1 2 3 4 5 6
Formic Acid (µmol)
Time (hours)
Figure 4.2 Time study preformed under standard reaction condition from 1-6 hours
56
Table 4.2 Effect of varying base
Entry Base
a
HCOOH (TON)
b
MeOH (TON)
b
HCHO (TON)
b
1 - 0.0 0.0 0.0
2 LiOH 1.4 0.0 0.0
3 NaOH 5.7 0.9 0.0
4 KOH 56.6 0.0 0.0
5 RbOH 15.5 0.6 0.0
6 CsOH 0.6 0.1 trace
7 Ca(OH)
2
0.2 0.5 0.0
8 AgTFA 0.0 0.2 0.5
9 Et
3
N 0.0 0.0 0.0
10 NH
4
OH 1.0 1.0 0.0
a)
1.7 mmol base used.
b)
Determined by wet1D NMR with a DMSO standard. Reaction
conditions: 5 µmol 1 and 1.7 mmol of base were dissolved in a 1.0 mL H
2
O. The reaction
mixture was then placed in a stainless steel reactor with a high-pressure valve composed
of Swagelok components and was purged and charged with 6.8 atm (806 µmol) carbon
dioxide gas, followed by 20 atm (2,370 µmol) hydrogen gas. The mixture was stirred at
200 °C for three hours.
4.2.3 Effects of Palladium Salts
With these reaction conditions in hand, we compared our palladium complex to
other known palladium(II) sources (Table 4.3). Critically, in the absence of a palladium
source no reaction took place (Table 4.3, entry 1). Palladium(II) chloride produced the
lowest yield of formic acid (Table 4.3, entry 3), while
bis(acetonitrile)dichloropalladium(II) produced the highest yield of the palladium salts
(Table 4.3, entry 4). Palladium(II) acetate was slightly less active, producing 43.1 TON
of formic acid (Table 4.3, entry 5).
57
Table 4.3 Effect of palladium catalysts
Entry Pd Catalyst HCOOH (TON)
a
1 - 0.0
2 1 56.6
3 PdCl
2
7.4
4 Pd(MeCN)
2
Cl
2
48.5
5 Pd(OAc)
2
43.1
a)
Determined by wet1D NMR with a DMSO standard. Reaction conditions: 5 µmol
palladium source and 1.7 mmol of KOH were dissolved in a 1.0 mL H
2
O. The reaction
mixture was then placed in a stainless steel reactor with a high-pressure valve composed
of Swagelok components and was purged and charged with 6.8 atm (806 µmol) carbon
dioxide gas, followed by 20 atm (2,370 µmol) hydrogen gas. The mixture was stirred at
200 °C for three hours.
4.2.4 Effects of Metals Salts in with NHC-amidate Ligand
In light of these impressive results of our NHC-amidate ligand, we continued to
pursue the versatility of our ligand by coupling it to various other metals (Table 4.4). In
order to do so, we prepared the catalyst system in situ and directly used the crude metal
complex. We initially tested the NHC-amidate ligand without the addition of a metal
(Table 4.4, entry 1). Surprisingly, 1.7 TON of formic acid was produced, which is
impressive since no product is formed in the absence of the ligand complex. We then
studied the hydrogenation of carbon dioxide with metals in the same group as palladium.
Though neither nickel(II) chloride nor the complexed system provided any significant
results (Table 4.4, entry 2), the reaction with platinum and the NHC-amidate ligand
provided impressive results (Table 4.4, entry 3). Bis(acetonitrile)dichloroplatinum
complexed with our NHC-amidate ligand produced 64 TON of formic acid, while only
1.3 TON of formic acid was observed when bis-(acetonitrile)dichloroplatinum was used
alone. Complexing our ligand with rhodium(III) chloride nearly doubled the results
58
compared to the metal salt alone (Table 4.4, entry 4), while the addition of silver by silver
oxide increased the yield 20-fold (Table 4.4, entry 5). Coupling the NHC-amidate ligand
with iridium and ruthenium metals had lesser affects, though higher yields were still
observed with the complexed systems (Table 4.4, entry 6-7). Lastly, our NHC-amidate
ligand was coupled with magnesium, using magnesium bromide as the metal source,
which again had higher yields than the sole metal in the reaction mixture (Table 4.4,
entry 8).
Table 4.4 Effect of varying metal salts and versatility of the NHC-amidate ligand
Entry
Metal Source HCOOH (TON)
b
M[OMe]
a
HCOOH (TON)
b
1 - - Ligand 1.7
2
NiCl
2
0.0
Ni[OMe] 0.0
3 (MeCN)
2
PtCl
2
1.3 Pt[OMe] 64.0
4
RhCl
3
3.8
Rh[OMe] 7.0
5
Ag
2
O 1.4
Ag[OMe] 28.6
6 IrCl
3
14.4 Ir[OMe] 17.6
7
RuCl
3
1.3
Ru[OMe] 3.1
8 MgBr
2
0.9 Mg[OMe] 4.4
a)
Varying metal complexes were prepared by stirring 20 µmol of the NHC-amidate
ligand was stirred with 20 µmol of the parent metal source in 1.0 mL aceotonitrile at
room temperature for 16 hours.
b)
Determined by wet1D NMR with a DMSO standard.
Reaction conditions: 5 µmol of metal source and 1.7 mmol of KOH were dissolved in a
1.0 mL H
2
O. The reaction mixture was then placed in a stainless steel reactor with a high-
pressure valve composed of Swagelok components and was purged and charged with 6.8
atm (806 µmol) carbon dioxide gas, followed by 20 atm (2,370 µmol) hydrogen gas. The
mixture was stirred at 200 °C for three hours.
4.2.5 Effects at Pressure and Heterogeneous Catalyst Conditions
In order to improve yields, higher pressures of carbon dioxide and hydrogen
where utilized (Table 4.5). Increasing the pressure of carbon dioxide to 68 atm and
59
hydrogen to 41 atm yielded 50% more formic acid than previous lower pressures with
trace amounts of methanol (Table 4.5, entry 1). We further emulated Takezaki et. al.
work by comparing our catalyst to palladium(II) chloride, which produced close to 140
TON of formic acid selectively (Table 4.5, entry 2). We then decide to employ a
heterogeneous catalyst system by excluding water from the reaction mixture. While the
results using palladium(II) chloride drastically decreased to 24 TON of formic acid
(Table 4.5, entry 3), our catalyst was able to produce 150 TON of formic acid selectively
in the absence of any solvent (Table 4.5, entry 4). This is in sharp contrast to not forming
any product when no catalyst was used (Table 4.5, entry 5).
Table 4.5 Effect of varying palladium sources and solvent amount
Entry
Pd Source CO
2
(psi) H
2
O (mL) HCOOH (TON)
MeOH (TON)
1
1 1,000
1.0 95.3 0.3
2
PdCl
2
1,000
1.0 139.7 0.0
3
PdCl
2
1,000
0.0 23.5 0.0
4
1 1,000
0.0 149.9 0.0
5
- 1,000
0.0 0.0 0.0
a
Reaction conditions: 1.7 mmol KOH and 5 µmol of a palladium source were add into a
stainless steel reactor and charged with 68 atm (8,000 µmol) carbon dioxide and 41 atm
(4,800 µmol) hydrogen gas. The mixture was heated at 200 °C for 3 hours.
b
Yields were
calculated using DMSO as an internal NMR standard.
4.2.6 Proposed Mechanism via a Palladium-Hydroxide Complex
Interestingly, in the absence of carbon dioxide, both potassium carbonate and
potassium bicarbonate produced significant amounts of formic acid. Further, we found
that in an aqueous alkaline media, carbon dioxide is readily converted into potassium
60
carbonate. Therefore, we believe the carbon source which reacts with catalyst system is
carbonate and not carbon dioxide (Scheme 4.2). Hydrogen gas forms a metal hydride
species with 1, which in turn reacts with carbonate. Carbonate is reduced to formate and
generates a Pd-OH species.
[9a, 13]
The subsequent release of a hydroxide ion regenerates
complex 1. In efforts to increase the amount of formate produced, we utilized known
carbonate forming agents;
[14]
however, the addition of Cu/ZnO or Mg(OH)
2
did not
increase the yield of the reaction.
Scheme 4.2 Proposed mechanism through a palladium-hydroxide complex
H
2
Pd
N
O
N
N
O
Me
Cl
HCOO
K
2
CO
3
Pd
N
OH
N
N
O
Me
Cl
OMe
OH
Pd
N
N
N
O
Me
Cl
OMe
H
H
61
4.2.7 Summary
In summary we have we have developed a set of reaction conditions which
selectively hydrogenates carbon dioxide gas into formic acid, which can be utilized in
alternative fuel sources, such as DFAFC. Notably, our NHC-amidate ligand can be
complexed with various metals other than palladium, and in many cases provide higher
yields of formic acid than simply using the metal salt without the ligand. In addition, by
making the reaction into a heterogeneous catalytic system, the highest yields of formic
acid were observed by utilizing our NHC-amidate Pd(II) catalyst. This is of great
importance in order to reduce the amount of greenhouse gasses and convert them into
possible alternative fuel sources.
4.3 Experimental Methods
To a 0.5 dram vial equipped with a stir bar was added a solution of 1.0 mL
distilled water, 95.3 mg (1.7 mmol) KOH, and 2.0 mg (5.0 µmol) complex 1. The vial
was then placed in a stainless steel reactor with a high-pressure valve composed of
Swagelok components and was purged and charged with 6.8 atm (806 µmol) carbon
dioxide gas, followed by 20 atm (2,370 µmol) hydrogen gas. The mixture was stirred at
200 °C for three hours. 0.5 mL D
2
O was added to the reaction mixture and a wet1D NMR
was taken using DMSO as an internal standard.
62
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73
Appendix 1
Supporting Information for Chapter 1
74
APPENDIX 1 TABLE OF CONTENTS
General Information………………...………………………………………..……….… 75
wet1D
1
H NMR Spectra,
12
C CH
4
, D
2
O workup (Table 1, entry 4)…….…………..….. 76
wet1D
1
H NMR Spectra,
13
C CH
4
, D
2
O workup (Figure 1)……….…………..……….. 77
wet1D
1
H NMR Spectra,
12
C CH
4
, TFA-d
1
workup…………………………...……….. 78
13
C NMR Spectra,
13
C CH
4
, D
2
O workup (Figure 2)………………………….……...... 79
13
C NMR Spectra,
12
C CH
4
, TFA-d
1
workup……………………………..…...……….. 80
19
F NMR Spectra, TFA-d
1
workup……………………………...…….…………….….. 81
wet1D
1
H NMR Spectra (Table 2, entry 2)…………………..…………….………….... 82
wet1D
1
H NMR Spectra (Table 2, entry 3)…………..…………..…………….……….. 83
wet1D
1
H NMR Spectra (Table 2, entry 4)……………...…………..…………....…….. 84
wet1D
1
H NMR Spectra (Table 2, entry 5)……….…………………..………......…….. 85
wet1D
1
H NMR Spectra (Table 3, entry 1)……………….………………..……..…….. 86
wet1D
1
H NMR Spectra (Table 3, entry 3)………………….………………………….. 87
wet1D
1
H NMR Spectra (Table 4, entry 1)…………...……………………..………….. 88
wet1D
1
H NMR Spectra (Table 4, entry 2)………….…………….…………...……….. 89
wet1D
1
H NMR Spectra (Table 4, entry 3)……………….………..……………..…….. 90
wet1D
1
H NMR Spectra (Table 4, entry 4)………………….…………..…………..….. 91
Gas Chromotography…………..……………………………………………..……….... 92
Density Functional Theory Calculations Supplemental Information…………………... 93
75
GENERAL INFORMATION
All glassware and reactor components were oven-dried prior to use. All chemicals
were purchased as reagent grade and used without further purification. TFA and TFAA
were purchased from Sigma Aldrich, hydrogen peroxide was purchased from Macron
Chemicals, and D
2
O was purchased from Cambridge Isotope Laboratories.
1
H (400 or
500 MHz),
13
C (126 MHz), and
19
F (470 MHz) NMR spectra were obtained on a Varian
Mercury 400 MHz or Varian VNMRS-500 and referenced to DMSO or TFA.
Thermo Finnigan Trace GC was used for the gas chromatography studies. TCD
detector was set at 160 °C and the inlet temperature was set at 200 °C. Initial oven
temperature was 40 °C for one minute and was then ramped to 50 °C at 10 °C/minute.
The flow rate was set to 3 mL/minute by an argon high purity carrier. The column used
was a Supelco Carboxen 1010, 30 meters x 0.53 mm. After reaction was complete, the
reactor was cooled at -20 °C for half an hour. The vessel was opened and the expelled gas
was passed through potassium hydroxide and Drierite and collected in a screw cap vial
with septum. Resulting gas was then injected into the GC column.
76
77
78
79
80
81
82
83
84
85
86
87
88
89
90
91
92
Gas chromatograph of standard functionalization reaction.
Conditions: 0.2 mL TFA, 0.6 mL TFAA, 13 µL (130 µmol) 30% H
2
O
2
, and 27 atm methane gas stirred at 60 °C for 16
hours.
93
DENSITY FUNCTIONAL THEORY CALCULATIONS SUPPLEMENTAL
INFORMATION
Format:
• Compound Name
• Electronic Energy (hartrees)
• Thermal Energy Correction (hartrees)
• Free Energy Correction (hartrees)
• Cartesian Coordinates for Optimized Geometry (Angstroms)
94
Trifluoromethyl radical (CF
3
•)
-337.6081691
-337.5934041
-337.6244541
C 0. 0. 0.32881
F 0. 1.26811 -0.07307
F -1.09822 -0.63405 -0.07307
F 1.09822 -0.63405 -0.07307
95
Methane (CH
4
)
-40.520797
-40.473319
-40.492924
C 0. 0. 0.
H 0.63333 0.63333 0.63333
H -0.63333 -0.63333 0.63333
H -0.63333 0.63333 -0.63333
H 0.63333 -0.63333 -0.63333
96
Transition State 1
-378.1115917
-378.0516267
-378.0927377
C 2.36837 -0.00004 0.00032
H 2.63789 -1.0407 0.20126
H 1.04149 0.00008 0.00044
H 2.63833 0.69426 0.80098
H 2.63796 0.34613 -1.00141
C -0.32889 -0.00002 0.00011
F -0.78472 1.24406 -0.2313
F -0.7845 -0.82219 -0.96192
F -0.78551 -0.42181 1.19278
97
Trifluoromethane (CF
3
H)
-338.28754
-338.271782
-338.30164
C 0. 0. 0.34365
H 0. 0. 1.44021
F 0. 1.26332 -0.12971
F -1.09407 -0.63166 -0.12971
F 1.09407 -0.63166 -0.12971
98
Methyl radical (CH
3
•)
-39.8449057
-39.8118187
-39.8332297
C 0. 0. 0.0005
H 0. 1.08777 -0.00099
H -0.94204 -0.54389 -0.00099
H 0.94204 -0.54389 -0.00099
99
Trifluoroacetic acid anhydride (TFAA)
-977.2802269
-977.2145479
-977.2741009
C -2.3581 -0.31219 0.05036
C -1.18571 0.66953 -0.21025
O -1.30936 1.79403 -0.57099
O -0.00001 -0.01128 0.00015
C 1.18576 0.66955 0.21032
C 2.35808 -0.31222 -0.0504
F 2.31216 -0.76392 -1.31977
F 2.2766 -1.36849 0.78253
F 3.52503 0.30444 0.14704
F -2.31232 -0.76388 1.31977
F -2.27656 -1.36849 -0.78251
F -3.525 0.30449 -0.14719
O 1.30945 1.79407 0.57097
100
Transition State 2
-1017.119449
-1017.017689
-1017.081752
C -2.48565 -0.33728 0.06395
C -1.28284 0.62469 -0.13337
O -1.39784 1.79313 -0.35225
O -0.13326 -0.08123 0.02684
F -2.43036 -0.91932 1.28071
F -2.46779 -1.30805 -0.87167
F -3.63701 0.33338 -0.03573
C 1.11077 0.4307 -0.44443
C 1.49788 1.96752 1.10701
H 2.52863 2.20002 0.8478
H 0.72345 2.65259 0.77012
H 1.31587 1.37637 2.00331
C 2.20109 -0.5801 0.01496
O 1.22998 1.09475 -1.44725
F 3.42281 -0.0535 -0.156
F 2.1069 -1.68747 -0.75068
F 2.0714 -0.95564 1.30131
101
Species 5
-1017.14548
-1017.040414
-1017.102606
C -2.46603 -0.34016 0.03045
C -1.26084 0.63475 -0.07317
O -1.3872 1.81665 -0.23789
O -0.12324 -0.06294 0.06168
F -2.46275 -0.96658 1.22777
F -2.40519 -1.27803 -0.93665
F -3.618 0.3271 -0.09474
C 1.16141 0.63431 -0.10958
C 1.40733 1.65928 1.02816
H 2.45087 1.98419 0.99684
H 0.74141 2.51234 0.8843
H 1.19056 1.15936 1.97907
C 2.18885 -0.55439 -0.00246
O 1.29437 1.22206 -1.29017
F 3.44191 -0.08148 -0.11837
F 1.98263 -1.45702 -0.97281
F 2.07935 -1.17341 1.18707
102
Transition State 3
-1017.128358
-1017.024944
-1017.089129
C -2.46763 -0.40764 0.03683
C -1.35813 0.67061 -0.09194
O -1.58949 1.83292 -0.28198
O -0.15662 0.09402 0.06254
F -2.41504 -0.99308 1.25464
F -2.31066 -1.3673 -0.89766
F -3.67732 0.14151 -0.11546
C 1.02198 0.95971 -0.15149
C 1.38357 1.72567 1.10271
H 0.64218 2.5317 1.19374
H 1.33625 1.09465 1.99505
H 2.37813 2.16675 0.98628
C 2.32211 -0.64502 -0.01169
O 1.27185 1.30757 -1.29563
F 3.54068 -0.13761 -0.11707
F 2.0504 -1.47113 -1.00026
F 2.14818 -1.22674 1.16904
103
Species 6
-679.5386106
-679.4523346
-679.5042376
C -1.58727 -0.18032 -0.05726
C -0.2331 0.50207 0.2732
O -0.13974 1.62143 0.67553
O 0.76964 -0.39713 0.07551
F -1.74868 -1.30356 0.67615
F -1.63133 -0.52167 -1.36352
F -2.60582 0.64496 0.20196
C 2.11241 0.04033 -0.14211
C 3.05973 -1.02541 0.31448
H 4.07855 -0.74915 0.02947
H 2.98754 -1.13126 1.40605
H 2.7807 -1.99005 -0.12887
O 2.35947 1.0848 -0.659
104
Appendix 2
Supporting Information for Chapter 2
105
APPENDIX 2 TABLE OF CONTENTS
General Information……………………………………….…….……………..……… 107
Styrene + cyclopentene
1
H NMR Spectra (Table 5, entry 1)…………………..……… 108
Styrene + cyclopentene
13
C NMR Spectra (Table 5, entry 1)……………………….... 109
Styrene + cyclopentene gCOSY Spectra (Table 5, entry 1)…………..………..…...… 110
Styrene + cyclopentene gHSQC Spectra (Table 5, entry 1)………………..…...…….. 111
Styrene + cyclopentene GC/MS (Table 5, entry 1)………..………..………..……….. 112
4-trifluoromethylstyrene + cyclopentene
1
H NMR Spectra (Table 5, entry 2)……...… 113
4-trifluoromethylstyrene + cyclopentene
13
C NMR Spectra (Table 5, entry 2)…...….. 114
4-trifluoromethylstyrene + cyclopentene GC/MS (Table 5, entry 2)…………...…….. 115
4-nitrostyrene + cyclopentene
1
H NMR Spectra (Table 5, entry 3)…………......….… 116
4-nitrostyrene + cyclopentene
13
C NMR Spectra (Table 5, entry 3)…………. ….…... 117
4-nitrostyrene + cyclopentene GC/MS (Table 5, entry 3)……….……………...…….. 118
4-chlorostyrene + cyclopentene
1
H NMR Spectra (Table 5, entry 4)……..…..…….… 119
4-chlorostyrene + cyclopentene
13
C NMR Spectra (Table 5, entry 4)………. ..…….... 120
4-chlorostyrene + cyclopentene GC/MS (Table 5, entry 4)……….………...…..…….. 121
2-fluorostyrene + cyclopentene
1
H NMR Spectra (Table 5, entry 5)…….. ………..… 122
2-fluorostyrene + cyclopentene
13
C NMR Spectra (Table 5, entry 5)….…….…... ….. 123
2-fluorostyrene + cyclopentene GC/MS (Table 5, entry 5)……….……….………….. 124
4-fluorostyrene + cyclopentene
1
H NMR Spectra (Table 5, entry 6)……….. .….…… 125
4-fluorostyrene + cyclopentene
13
C NMR Spectra (Table 5, entry 6)………..... …….. 126
4-fluorostyrene + cyclopentene GC/MS (Table 5, entry 6)……….…….…………….. 127
4-methylstyrene + cyclopentene
1
H NMR Spectra (Table 5, entry 7)…..….... ………. 128
4-methylstyrene + cyclopentene
13
C NMR Spectra (Table 5, entry 7)………..….. ….. 129
4-methylstyrene + cyclopentene GC/MS (Table 5, entry 7)……….……..……..…….. 130
2-vinylnaphthalene + cyclopentene
1
H NMR Spectra (Table 5, entry 8)…...………… 131
2- vinylnaphthalene + cyclopentene
13
C NMR Spectra (Table 5, entry 8)……..…... ... 132
2- vinylnaphthalene + cyclopentene GC/MS (Table 5, entry 8)…………….…….…... 133
4-methoxystyrene + cyclopentene
1
H NMR Spectra (Table 5, entry 9)……..……… 134
106
4-methoxystyrene + cyclopentene
13
C NMR Spectra (Table 5, entry 9)…………….... 135
4-methoxystyrene + cyclopentene GC/MS (Table 5, entry 9)……….……….……….. 136
107
GENERAL INFORMATION
All glassware were oven-dried prior to use. All chemicals were purchased as
reagent grade and used without further purification.
1
H (400 or 500 MHz) and
13
C (126
MHz) NMR spectra were obtained on a Varian Mercury 400 MHz or Varian VNMRS-
500 and referenced to DMSO.
108
109
110
111
112
113
114
115
116
117
118
119
120
121
122
123
124
125
126
127
128
129
130
131
132
133
134
135
136
137
Appendix 3
Supporting Information for Chapter 3
138
APPENDIX 3 TABLE OF CONTENTS
DMSO Calibration Curve with Formic Acid………………...……..……………….… 139
DMSO Calibration Curve with Glycolic Acid…………………………..…..………… 140
Lewis Acid Additive Results………..……………………………..………..………… 141
D-Glucose wet1D NMR Spectra (Table 2, entry 3)…………………...……..……….. 142
D-Glucose
13
C NMR Spectra (Table 2, entry 3)………………..….………………….. 143
D-Galactose wet1D NMR Spectra (Table 4, entry 1)…………………..……….…….. 144
D-Ribose wet1D NMR Spectra (Table 4, entry 2)…………………….……..……….. 145
D-Xylose wet1D NMR Spectra (Table 4, entry 3)……….………………..………….. 146
D-Fructose wet1D NMR Spectra (Table 4, entry 4)………………..….…...…...…….. 147
D-Tagatose wet1D NMR Spectra (Table 4, entry 5)…………...…………..….…….... 148
D-Maltose wet1D NMR Spectra (Table 4, entry 6)…………………………..……….. 149
D-Lactose wet1D NMR Spectra (Table 4, entry 7)……………………...…..….…….. 150
D-Cellobiose wet1D NMR Spectra (Table 4, entry 8)……….…...……….....……….. 151
Sucrose wet1D NMR Spectra (Table 5, entry 1)……………………..……………….. 152
D-Melezitose wet1D NMR Spectra (Table 5, entry 2)………….…………………….. 153
D-Raffinose wet1D NMR Spectra (Table 5, entry 3)………..………………....….….. 154
Glycerol wet1D NMR Spectra (Table 5, entry 4)……….…………………………….. 155
139
DMSO Calibration Curve with Formic Acid
The calibration curve of a DMSO standard was taken with known amounts of formic
acid. The
1
H NMR of 20, 40, 80, and 120 µmol of formic acid was taken in 0.75 mL of
D
2
O with a sealed capillary DMSO standard.
y = 10.315x
R² = 0.99372
0
20
40
60
80
100
120
140
0 2 4 6 8 10 12 14
micromol of Formic Acid
NMR Integration
140
DMSO Calibration Curve with Glycolic Acid
The calibration curve of a DMSO standard was taken with known amounts of glycolic
acid. The
1
H NMR of 1, 5, 10, and 20 µmol of glycolic acid was taken in 0.75 mL of D
2
O
with a sealed capillary DMSO standard.
y = 3.4771x
R² = 0.99756
0
5
10
15
20
25
0 1 2 3 4 5 6 7
micromol of Glycolic Acid
NMR Integration
141
Lewis Acid Additive Results
Lewis Acid TON
AlCl
3
34.6
CrCl
3
79.9
ZnCl
2
64.9
SnCl
2
51.5
Reaction conditions: 100 µmol of substrate, 5 µmol of Lewis acid, and 600 µmol NaOH
were dissolved in 0.44 mL H
2
O. 60 µL 30% H
2
O
2
was added and the mixture stirred at
25 °C for 16 hours. 0.25 mL of D
2
O was then added to the reaction mixture with a sealed
capillary DMSO standard. The solution was then analyzed using wet1D NMR.
142
143
144
145
146
147
148
149
150
151
152
153
154
155
156
Appendix 4
Supporting Information for Chapter 4
157
APPENDIX 4 TABLE OF CONTENTS
General Information…………………………………………………………………… 158
Standard reaction condition wet1D NMR Spectra (Table 1, entry 3)…………….…… 159
Standard reaction condition
13
C NMR Spectra (Table 1, entry 3)…………………….. 160
Heterolytic reaction condition wet1D NMR Spectra (Table 4, entry 4)……..………... 161
158
GENERAL INFORMATION
All glassware were oven-dried prior to use. All chemicals were purchased as
reagent grade and used without further purification.
1
H (400 or 500 MHz) and
13
C (126
MHz) NMR spectra were obtained on a Varian Mercury 400 MHz or Varian VNMRS-
500 and referenced to DMSO.
159
160
161
Abstract (if available)
Abstract
Due to the composition of organic compounds, the activation of carbon and hydrogen bonds has grasped the attention of many modern chemists. The focus of the following projects was centered around the activation and better understanding of C-H bonds. From mechanistic studies of greenhouse gasses, such as methane and carbon dioxide, co-dimerization of small molecules, or degradation of abundant biomass, the breakage of C-H bonds and subsequent reformation of new C-H bonds are of paramount importance. ❧ In Chapter 1, thorough mechanistic studies and density functional theory (DFT) calculations revealed a background radical pathway latent in metal catalyzed oxidations of methane. Use of hydrogen peroxide with trifluoroacetic anhydride (TFAA) generated a trifluoromethyl radical (·CF₃), which in turn reacted with methane gas selectively to yield acetic acid. It was found that the methyl carbon of the product was derived from methane and the carbonyl carbon was derived from TFAA. Computational studies also support these findings, revealing each step of the cyclic reaction is energetically favorable. ❧ A highly efficient co-dimerization of styrene and cyclopentene, in the presence of palladium and a BF₃ source, was developed selectively forming a new C-C bond, in Chapter 2. The complex [Pd(PPh₃)₂]⁺ BF₄⁻ is believed to generate palladium hydride (Pd-H), which catalyzes the reaction between various styrenes and cyclopentene in excellent yields as single isomers. This co-catalytic system provides a new, efficient C-C bond forming method. ❧ The development of an effective method that converts a variety of mono- and disaccharides into predominantly formic acid is discussed in Chapter 3. A recyclable NHC-amidate palladium(II) catalyst facilitates oxidative degradation of carbohydrates without using excess oxidant. Stoichiometric amounts of hydrogen peroxide and sodium hydroxide were employed at ambient temperatures. ❧ Lastly, Chapter 4 discusses a highly efficient hydrogenation of carbon dioxide into formic acid in the presence of the same NHC-amidate palladium(II) complex. Excellent turnover number is observed when the catalyst was used under heterolytic conditions. This catalytic system provides a new efficient carbon dioxide hydrogenation method.
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Zargari, Nima
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Carbon-hydrogen bond activation: radical methane functionalization; unactivated alkene coupling; saccharide degradation; and carbon dioxide hydrogenation
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College of Letters, Arts and Sciences
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Doctor of Philosophy
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Chemistry
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04/19/2016
Defense Date
03/07/2016
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Tag
alkene cross coupling,biomass,carbon dioxide,hydroalkenylation,hydrogen peroxide,hydrogenation,methane functionalization,OAI-PMH Harvest,palladium catalysis,Pd-hydride,radical,recyclable catalyst,tetrafluoroborate,trifluoroacetic anhydride,trifluoromethyl radical
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Tags
alkene cross coupling
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methane functionalization
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Pd-hydride
radical
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tetrafluoroborate
trifluoroacetic anhydride
trifluoromethyl radical